JOHN ALEXANDER JAMESON, JR. 1903-1934 ENGINEERING LIBRARY THIS BOOK belonged to John Alexander Jameson, Jr., A.B., Wil- liams, 1925; B.S., Massachusetts Institute of Technology, 1928; M.S., California, 1933. He was a member of Phi Beta Kappa, Tau Beta Pi, the American Society of Civil Engineers, and the Sigma Phi Fraternity. His untimely death cut short a promising career. He was engaged, as Research Assistant in Mechanical Engineering, upon the design and construction of the U. S. Tidal Model Labora- tory of the University of California. His genial nature and unostentatious effectiveness were founded on integrity, loyalty, and devotion. These qualities, recognized by everyone, make his life a continuing beneficence. Memory of him will not fail among those who knew him. ' INTERNATIONAL CHEMICAL SERIES H. P. TALBOT, PH.D., Sc.D., CONSULTING EDITOR A TEXTBOOK OF INORGANIC CHEMISTRY FOR COLLEGES INTERNATIONAL CHEMICAL SERIES (H. P. TALBOT, PH.D., Sc.D., CONSULTING EDITOR) Bancroft APPLIED COLLOID CHEM- ISTRY Bingha m FLUIDITY AND PLASTICITY Cady INORGANIC CHEMISTRY Cady GENERAL CHEMISTRY Griffin TECHNICAL METHODS OF ANALYSIS As Employed in the Labora- tories of Arthur D. Little, Inc. Hall and W illiams CHEMICAL AND METALLO- GRAPHIC EXAMINATION OF IRON, STEEL AND BRASS Hamilton and Simpson CALCULATIONS OF QUAN- TITATIVE ANALYSIS Loeb PROTEINS AND THE THEORY OF COLLOIDAL BEHAVIOR Lord and Demorest METALLURGICAL ANALY- SIS Fifth Edition Mahin QUANTITATIVE ANALYSIS Third Edition Mahin and Carr QUANTITATIVE AGRICUL- TURAL ANALYSIS Millard PHYSICAL CHEMISTRY FOR COLLEGES CHEMICAL Moore HISTORY OF CHEMISTRY Norris TEXTBOOK OF INORGANIC CHEMISTRY FOR COL- LEGES Norris and Mark LABORATORY EXERCISES IN INORGANIC CHEMIS- TRY Norris ORGANIC CHEMISTRY Second Edition Norris EXPERIMENTAL ORGANIC CHEMISTRY Second Edition Parr ANALYSIS OF FUEL, GAS. WATER AND LUBRICANTS Third Edition Robinson THE ELEMENTS OF FRAC- TIONAL DISTILLATION White TECHNICAL GAS AND FUEL ANALYSIS Second Edition Williams PRINCIPLES OF METALLO- GRAPHY Woodman FOOD ANALYSIS Second Edition Long and Anderson CHEMICAL CALCULATIONS BoQue THE THEORY AND APPLI- CATION OF COLLOIDAL BEHAVIOR Two Volumes A TEXTBOOK OF INORGANIC CHEMISTRY FOR COLLEGES BY JAMES F. NORRIS PROFESSOR OF ORGANIC CHEMISTRY MASSACHUSETTS INSTITUTE OF TECHNOLOGY AUTHOR OF "THE PRINCIPLES OF ORGANIC CHEMISTRY," "EXPERIMENTAL ORGANIC CHEMISTRY," AND (WITH K. L. MARK) "LABORATORY EXERCISES IN INORGANIC CHEMISTRY" FIRST EDITION FIFTH IMPRESSION McGRAW-HILL BOOK COMPANY, INC. NEW YORK: 370 SEVENTH AVENUE LONDON: 6 & 8 BOUVERIE ST., E. C. 4 1921 Xt, COPYRIGHT, 1921, BY THE MCGRAW-HILL BOOK COMPANY, INC. PRINTKQ.IN^THE UNITED, STATES OF AMERICA PREFACE FOR a number of years the author of this textbook had the opportunity to teach students who were beginning the study of chemistry. His experience, acquired in the recitation room and in the laboratory, led him to the view that the average student finds it difficult to understand many of the apparently simple concepts of the science. It appeared, therefore, to be an interest- ing task to attempt to present the material commonly treated in elementary books on chemistry in a form which could be reason- ably well followed by the student through private study and with the smallest amount of explanation on the part of the teacher. Since this book has been written from this point of view, the sub- ject has been developed slowly, and the consideration of the more abstruse material has been deferred until the student has gained some familiarity with chemical phenomena and with the language of the science. No attempt has been made at conciseness in the discussion of important principles. Analogies have been repeat- edly pointed out in an endeavor to indicate to the student the way in which he should classify the facts brought to his attention. The aim of the author has been to present the general princi- ples underlying the science; as a consequence, chemical phe- nomena have been discussed from the standpoints of both matter and energy. The law of mobile equilibrium in its broadest sense has been used repeatedly in interpreting many important facts. The more elementary parts of thermochemistry and electrochemis- try have also been emphasized. In fact, physical chemistry has been drawn on frequently, but an endeavor has been made to limit its use to the elucidation of the more important facts of inorganic chemistry. Several chapters of the book are devoted to the consideration, in a general way, of the physical and chemical properties of metals, non-metals, acids, bases, and salts. It is possible by this pro- vi PREFACE cedure to bring out generalizations of value and to emphasize the relations that exist between the uses of compounds and their properties. Important descriptive matter has not been omitted, and the more recent advances in chemistry in both its technical and theoretical aspects have been included. The modern con- ception of the atom has been presented in an elementary way. Throughout the book an attempt has been made to impress upon the student the fact that chemistry is a growing science. A large number of exercises are placed at the ends of most of the chapters. In the main, these are not direct questions on the text, but are designed to furnish the student with an opportunity to use his knowledge. Some of the questions may tax the abil- ity of even the best students; these may serve as the basis for discussions in the classroom. The author is greatly indebted to his wife for the efficient help which he has received throughout all the work required in the preparation of the book. This assistance included the copy- ing of the manuscript and help in proofreading and in the prepara- tion of the index. He is also indebted to Professor K. L. Mark, of Simmons College, for helpful suggestions and criticisms based on a careful reading of the manuscript. Four of the cuts have been borrowed, with the permission of the publishers, from Cady's Inorganic Chemistry. These include the excellent drawings to represent the chamber process for sul- phuric acid and the changes which take place in a blast furnace. Other cuts have been adapted from Leighou's Chemistry of Materials and Black and Conant's Practical Chemistry. JAMES F. NORMS. CAMBRIDGE, MASS. May 19, 1921. CONTENTS CHAPTER PAGE I. INTRODUCTION 1 II. PHYSICAL AND CHEMICAL CHANGES 6 III. ELEMENTS AND COMPOUNDS 15 IV. OXYGEN 21 V. HYDROGEN 38 VI. THE ATOMIC THEORY. CHEMICAL EQUATIONS 52 VII. CHEMICAL CALCULATIONS 68 VIII. MEASUREMENT OF GASES 77 IX. WATER 88 X. CHLORINE. VALENCE 99 XI. HYDROCHLORIC ACID. DOUBLE DECOMPOSITION 123 XII. THE ENERGY FACTOR IN CHEMICAL CHANGE 137 XIII. OZONE AND HYDROGEN PEROXIDE 144 XIV. PROPERTIES OF GASES, LIQUIDS, AND SOLIDS 155 XV. CARBON AND ITS OXIDES 171 XVI. COAL, COKE, ILLUMINATING GAS, FLAMES 194 XVII. ACIDS, BASES, SALTS. SOLUTIONS 209 XVIII. CHEMICAL EQUILIBRIUM 233 XIX. SULPHUR AND HYDROGEN SULPHIDE 245 XX. THE OXIDES AND ACIDS OF SULPHUR 256 XXI. NITROGEN AND THE ATMOSPHERE 284 XXII. AMMONIA AND ITS DERIVATIVES 302 XXIII. NITRIC ACID, NITROUS ACID, AND THE OXIDES OF NITROGEN. 320 XXIV. THE DETERMINATION OF ATOMIC AND MOLECULAR WEIGHTS. 345 XXV. THE PERIODIC LAW 358 XXVI. THE HALOGEN FAMILY 364 XXVII. SELENIUM AND TELLURIUM 392 XXVIII. PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 396 XXIX. SOME IMPORTANT ORGANIC COMPOUNDS 421 XXX. SILICON AND BORON. THE ACID-FORMING ELEMENTS AND THE PERIODIC CLASSIFICATION : 428 XXXI. THE PHYSICAL PROPERTIES OF METALS. ALLOYS 442 XXXII. THE CHEMICAL PROPERTIES OF METALS. METALLURGY... 457 XXXIII. ELECTROCHEMISTRY 471 XXXIV. THE PROPERTIES OF OXIDES, HYDROXIDES, AND SALTS.... 499 XXXV. SODIUM, POTASSIUM, RUBIDIUM, AND CAESIUM 515 XXXVI. CALCIUM, STRONTIUM, BARIUM. AND RADIUM 531 vii viii CONTENTS CHAPTER PAGE XXXVII. BERYLLIUM, MAGNESIUM, ZINC, CADMIUM, AND MERCURY. 551 XXXVIII. ALUMINIUM 566 XXXIX. TIN AND LEAD 578 XL. COPPER, SILVER, AND GOLD 588 XLI. IRON, COBALT, AND NICKEL 603 XLII. THE PLATINUM METALS 623 XLIII. CHROMIUM, MOLYBDENUM, TUNGSTEN, AND URANIUM 626 XLIV. MANGANESE 639 XLV. RADIOACTIVITY. THE STRUCTURE OF ATOMS 646 APPENDIX. . 659 INORGANIC CHEMISTRY CHAPTER I INTRODUCTION 1. What is chemistry, and why should the science be studied by one who is seeking a liberal education? If time of the greatest value is to be given to the study of -& subject, there should be an appreciation at the outset of what is to be gained by the effort. Will the knowledge acquired and the logical - inpde of, thinking developed be worth the labor? Shall we know the things about us better and thus increase the pleasure of living? Let us first see what chemistry is. Countless changes are taking place in the material world around us. In many of these, substances are undergoing pro- found transformations. Wood and coal burn and are changed into ashes; iron rusts and falls to a powder; sweet milk becomes sour; a seed sprouts, and feeding on the substances in the soil and in the air transforms them into a luxuriant plant; and the plant dies and becomes but a part of the soil again. The chemist studies such changes as these and records the properties of the substances which undergo transformation and the properties of the substances formed; he discovers under what circumstances such changes take place, and, with this knowledge, is able to direct many of them at will. He does more; he formulates laws which express the general behavior of substances, and, guided by these laws, brings about countless changes which never took place before. For centuries men with inquiring minds have observed nature closely, and have recorded and brought into systematic relation- ship many transformations in the substances which make up the 2 INORGANIC CHEMISTRY FOR COLLEGES material world. This knowledge forms the basis for the facts and laws of chemistry. The science is thus a vast storehouse of infor- mation of the greatest interest and value. It tells us, for example, what happens when silver tarnishes, when plants grow, and when bread turns sour. But more than this it teaches us how to prevent the metal from tarnishing, or how to remove the stain once it is formed; it shows us what to add to the soil to facilitate the growth of the plant, and how to make better bread, without the risk of the formation of the undesired acids. 2. The study of the changes which take place in nature of themselves, is but a small part of the science of chemistry. New substances of the greatest value have been made by the chemist, and ways have been found to produce new and desirable effects. The chemist has found out how to convert iron into a steel which possesses the properties required in tool making; he has made a steel for armor plate designed to resist most effectively the impact of projectiles; he has produced a third variety with properties which adap'i*/ it to the requirements exacted in the building of light but strong automobile parts. The chemist has changed agriculture from an art to a science, and has converted disagree- able waste products into valuable plant foods. He has made available processes in photography which place it among the fine arts; he has furnished antiseptics which prevent decay and dis- ease. The chemist builds electric batteries, bleaches cotton and wool, makes cements which displace wood and stone as building materials, and transforms the tar obtained by heating coal into dyes of every hue, into medicines, and into explosives of the greatest power. We truly live in the age of chemistry. Chemistry plays such an important part in our life every day, that we are constantly in contact with it. Life itself is associated with changes in the materials of which the body is made up. The digestion and assimilation of foods are chemical processes of the profoundest interest. The clothes we wear, and nearly every- thing which adds to our comfort and pleasure in life, have been under the hands of the chemist who has prepared them for our use. 3. Will not a knowledge of this broad and wonderful science add to our mental satisfaction and increase the pleasure of living? Have we not advanced in our education when we have learned to INTRODUCTION 3 interpret changes of the greatest interest which are taking place all around us? But the accumulation of useful knowledge is not the sole benefit gained from a study of chemistry. Other results of the greatest educational value follow, when the study of the science is conducted in an intelligent manner. In the laboratory the student examines the properties of typical and important substances, and investigates the transformations of these sub- stances into others possessing different properties. He thus has an opportunity to observe closely and to record his observations. He learns to see things as they are a power which comes only with training. The comprehension and the use of the laws and theories of chemistry develop the power of reasoning. The study of cause and effect and the relation between fact and theory, produce logical methods of thought which are characteristic of the educated man. Careful and conscientious work on the part of the student in meeting the requirements in a course in science will yield many by-products of the greatest value. He will develop promptness, accuracy, self-reliance, and clear methods of thought and expression. In the case of students who do not continue the study of chemistry beyond the first year, these by-products prove of greater value than the knowledge of facts acquired. Many of the latter may be soon forgotten, but the mental habits formed persist. 4. Birth of Chemistry. From the earliest times men have observed the more striking changes that take place in nature and have made accidental discoveries of great value. The stone age was followed by the age of bronze when the way to work metals was found out. Many of the substances used to-day, which are prepared by chemical methods, were known to the ancients, but there are no historical records of their discovery. Glass, for example, which has been known for over 3000 years, was prob- ably first made in China. According to tradition a method of making glass was discovered independently by some Phoenician sailors, who were preparing their food on a sandy beach. They supported their cooking utensils on pieces of crude soda taken from the cargo of their vessel. The heat of the fire brought about a reaction between the sand and soda, and a transparent glassy substance was formed. According to another story the dis- covery was made by Egyptian priests in connection with the 4 INORGANIC CHEMISTRY FOR COLLEGES extraction of gold from the earthy material with which it was mixed. They added soda to the mixture to lower the tempera- ture to which the mass was heated to melt it, and found that a transparent substance was formed. Soap was also made in the earliest times; wood ashes and fat were used and it is easy to see how the process could be discovered by accident. Important discoveries were made by the Egyptian priests, but they were kept as secrets, and many of them were forgotten and redis- covered later. But advance was slow as long as observation and chance alone led the way. Directed experimentation was necessary; and for this a motive had to be present. It was with the rise of alchemy that such a motive appeared. The alchemists sought the things which were supposed to lead to happiness health and riches. They endeavored to change the common metals into gold, hoping to do this with the help of a mysterious substance, called the philosopher's stone. But gold without health is of little value, so a search was made for the elixir of life, which could bring back glorious youth to the aged. Men worked at these problems all over Europe. They studied everything available, mixed things together, and heated and distilled them when possible. As sub- stances appeared to affect one another more readily when they were dissolved in some liquid, a third substance, a universal solvent, was sought as an aid toward accomplishing the end in view. Many important discoveries were made as the result of the eager search by the alchemist for what proved to be unattain- able, and some of the processes used to-day in chemistry were invented. The alchemists refer repeatedly to distillation, extrac- tion, calcination, coagulation, etc. They prepared and studied many of the compounds which are to-day commonly used in chemical work. Among these are sulphuric, nitric, and hydro- chloric acids, alum, soda, ammonium chloride, niter, and com- pounds of mercury, arsenic, and antimony. Some of the alchem- ists gained great fame as a result of their work and writings. Geber, an Arab who lived in the eighth century, exerted an influ- ence on the later alchemists of the Middle Ages. His books were translated into Latin in the thirteenth century and were the standard for his followers. Other noted alchemists were Albertus INTRODUCTION 5 Magnus (1205-1280), Roger Bacon (1214-1294), and Raymond Lullus (1235-1315). Alchemy reached its height in the thirteenth to fifteenth centuries. It was customary for kings and princes to have alchemists attached to their courts, for the rewards would be great if success crowned their efforts. There was accumulated in this way a vast amount of information upon which the science of chemistry was later founded. The knowledge spread slowly, as each alchemist guarded his secrets carefully. In order to gain renown for his accomplishments, and thus, per- haps, win a place at court, he had to publish to the world his discoveries. But he must keep his secrets. As a result, alle- gorical and figurative language was used in such a way that both aims were reached; the discoveries were so described that they appeared of the greatest importance, but at the same time the description was written in more or less unintelligible language. The selfish and utilitarian motive which guided the alchemist could not lead to the development of a science. It was only when a desire on the part of men to learn the truth about the wonders of nature became the incentive, that chemistry as a science was born. CHAPTER II PHYSICAL AND CHEMICAL CHANGES 5. We have already learned that chemistry is a science that deals with those changes in nature in which the substances before and after the change are different. When iron is converted into iron rust we have a chemical change; for the powder formed is evidently a different substance from the hard and rigid metal from which it was produced. As it is of the first importance to have at the outset as clear a conception as possible of what are classed as chemical phenomena, a few simple experiments will be carefully studied. 6. Physical and Chemical Changes. Let us consider first what happens when two metals, platinum and magnesium, are heated in a flame. A wire of platinum when heated becomes hot and gives off light we could see it in a dark room. A change has taken place in it, for when hot it is red, whereas when cold it has the color of silver. If the wire is removed from the flame and allowed to cool, it will be found that the metal has all the properties it possessed before being heated the substance has not been converted into a new substance as a result of being heated. If we carry out a similar experiment with magnesium, the result is quite different. The metal burns with an intense white light, and is changed into a white powder a new substance is formed. This experiment with the two metals is very instruct- ive, for it furnishes examples of the two kinds of changes which may take place in the things around us. In the case of the plati- num the change did not yield a new substance; it is an example of what is called a physical change. In the case of the mag- nesium, a new substance was formed and the change was chemical. The effect of heat on platinum as the result of which the metal gives off light, softens, expands, etc., is studied in the science of physics. The change of metallic magnesium, when heated, into 6 PHYSICAL AND CHEMICAL CHANGES 7 the white powder, which is called magnesium oxide, is studied in the science of chemistry. 7. In the experiment just described the results were brought about by the application of heat. Other agencies produce physi- cal and chemical changes. Two experiments will illustrate how so-called mechanical energy can produce these effects. If a stick of sulphur is brought into contact with a pith-ball suspended by means of a silken thread, the substances have no effect on each other. If, now, the sulphur is rubbed briskly with a piece of silk or with a cat's skin, and again brought into contact with the ball, it will be found that the two substances attract each other, and the ball clings to the sulphur. If the two are now separated they repel each other; when the sulphur is brought near the pith-ball the latter moves away in order to get as far off from the sulphur as possible. The mechanical energy exerted on the sulphur has changed it; electricity has been produced on its surface, and this brings about the difference in its action on the pith-ball. But it is still sulphur. The change produced is a physical one and is studied in detail in a course in physics. . . When many substances are rubbed together more profound changes take place than that which was the result of the experi- ment just described. Mercury and iodine furnish an excellent example. Mercury is a heavy liquid, with the color of silver, and is recognized as the substance used in thermometers. Iodine is a black solid, which dissolves in alcohol, forming a brown liquid, known as tincture of iodine. When the two substances are rubbed together, in the presence of a little alcohol, which makes the change take place more rapidly, they are transformed into a red powder. The mercury and the iodine disappear and a substance with new properties, called mercuric iodide, is produced. The mechanical energy expended served to bring the two substances into intimate contact, and a chemical change took place. 8. Light is a form of energy which produces changes. Certain dyes fade in sunlight; the art of photography is based upon the action of light on substances containing silver and other metals; and the growth of plants requires this form of energy. Two simple experiments will be instructive. A radiometer, Fig. 1, consists of two rectangular pieces of thin sheet metal set at right angles and suspended at the point of intersection in such a way 8 INORGANIC CHEMISTRY FOR COLLEGES that the whole can rotate freely. Alternate sides of the plates are blackened. This arrangement is put into a glass vessel from which the air is exhausted, in order to reduce to a minimum the friction produced when the plates rotate. When a radiometer is placed in sunlight the vanes rotate rapidly. The light falling on the blackened surfaces repels them and thus produces the rotation. The effect is produced by radient energy. If the heat radiations are absorbed by passing the light through a proper screen, the rotations continue. Light causes in this case a phys- ical change. If a piece of paper is dipped into a solution made by dissolving silver nitrate in water, and is exposed to light, it will turn black. Silver nitrate is made by heating metallic silver with nitric acid; it is a white crystalline substance. When exposed to sunlight, in contact with paper, it is decomposed and the metal is produced in a form which is black. A new substance is the result of the change; silver nitrate is converted into silver. Light in this case effects a chemical change. 9. Electricity produces physical and chemical changes. If a wire through which a current of electricity is passing is brought over a compass needle, the latter will be deflected and will no longer point to the north. When the wire is removed the needle returns to its original position. The electricity produced, evidently, a change which was physical. This phenomenon is of the greatest importance, and has been studied in detail; upon it is based the transformation of electricity into mechanical power as exemplified in the electric motor. A chemical change can be readily produced by means of elec- tricity. If the ends of two platinum wires connected with a source of electricity are placed in water which contains a little acid, bubbles of gas will be seen to form on the wires in the solu- tion, and to rise to the surface. Water does not allow electricity to flow through it readily and the small amount of acid is added te make it possible for the current to flow. If the electricity is allowed to pass through the solution for a long enough time, the PHYSICAL AND CHEMICAL CHANGES 9 water will gradually disappear and large amounts of the gases will be produced. The change produced is, evidently, a chemical one; the water is converted into other substances, which possess entirely different properties from it. A piece of apparatus which is commonly used to demonstrate this experiment is shown in Fig. 2. It was devised by the chemist Hofmann and is known by his name. The ends of the wire connected with the source of electricity are joined to the apparatus at A and B. At these points wires pass through the glass and terminate in platinum plates, which are called the electrodes. The gas that rises from the electrodes collects in the two tubes and forces the water up into the receiver C. The gases formed can be removed from the apparatus and collected by opening the stop-cocks at the ends of the tubes. The Hofmann apparatus is a very convenient one with which to study decompositions yielding gases produced by passing electricity through solutions. We shall have occasion to refer to it repeatedly. Of the four kinds of energy which bring about physical and chemical change, heat and electricity are used extensively in chemistry. The application of light is limited at present largely to photographic processes. Mechani- cal energy is used to some extent when chemical changes between gases are brought about, for when the latter are highly compressed the results are, in cer- tain cases, more satisfactory than under ordinary conditions. 10. Variable and Characteristic Properties. There is much that can be learned from a consideration of the experiments just described. We have seen that the changes brought about in matter can be divided into two classes: those in which a sub- stance is converted into another substance a chemical change, and those in which no new substance appears a physical change. In order to determine to which class any given change belongs, it is necessary to note carefully the properties of all materials involved before and after the change. We recognize the different forms of matter by their properties. The common ones made FIG. 2. 10 INORGANIC CHEMISTRY FOR COLLEGES use of are those which appeal directly to our senses. Through this means we recognize form, luster, solubility, and color; hard- ness, brittleness, and ductility; odor and taste. Examine a crys- tal of sugar a piece of rock candy and see how it can be identi- fied by its properties. It has a definite shape and shiny surfaces, is transparent, colorless, and soluble in water, is hard and brittle, is odorless, and possesses a sweet taste. The identity of this particular piece of sugar is determined by all these properties; we could recognize it among other things and other pieces of sugar. But must all samples of sugar possess all of these properties? Suppose we grind the crystal to a fine powder. The form has changed, it is not transparent, and we cannot see shiny surfaces; it no longer possesses many properties of the crystal. Is it still sugar? Has a new substance been formed? Was the change a physical or chemical one? The powder has two properties which the crystal possessed; it dissolves in water and the solution has a sweet taste. We know it is sugar; and the change is a physical one. It is evident, therefore, that we must differentiate two kinds of properties those which serve to identify any particular piece of matter, which is called a body, and those which are associated with the material of which the body is made up. Properties of the former class are called variable properties, and those of the latter, characteristic The shape of a piece of sugar is a vari- able property; its sweet taste when dissolved in water is a characteristic property. When chemical changes take place, sub- stances are formed which possess different characteristic properties from those of the substances entering into the change. A difference in a single property of this type shows that another substance has been formed. In enumerating characteristic properties it is necessary to state the external conditions to which the substance is subjected. A characteristic property of water is that it is a liquid; but this statement is found to be true only when water is examined within a certain range of temperature; below the freezing-point it becomes a solid. Referring back to our experiment with the platinum wire, it will be remembered that the metal possessed the color of silver before being heated, and that when the change was brought about this disappeared and the wire became red. Was the change physical or chemical? The most evident change in property was PHYSICAL AND CHEMICAL CHANGES 11 in color. We must compare properties under the same external conditions; so we let the wire cool and examine it. It has the color of silver and has the other characteristic properties it pos- sessed before being heated; the change is thus a physical one. 11. Physical Properties and Chemical Properties. The properties which substances possess are classified, at times, from a point of view different from that used in dividing them into accidental and characteristic. A distinction is drawn in this case between the properties which we recognize directly, or when a physical change takes place, and those which become evident only when a chemical change occurs; the former are called physi- cal properties, the latter chemical properties. Some examples will make this distinction clear. Iron is hard, can be drawn out into a wire, is heavy, and allows electricity to pass through it; it has, thus, the properties of hardness, ductility, density, and elec- trical conductivity. All these properties can be recognized with- out a chemical change taking place in the metal; they are physical properties. Iron rusts when exposed to the air, and dissolves when put into an acid, such as vinegar, for example. These properties become evident only when the iron undergoes a chem- ical change; they are chemical properties. Since we recognize substances by both physical and chemical properties, attention must be paid to the two classes in the study of chemistry. The distinction which has just been drawn is often expressed in a different way. Physical properties are classed simply as properties, and what we have called the chemical properties of a substance are -referred to as its chemical conduct. The so-called chemical properties, as has been said, become evident only when the substance possessing them enters into chemical reaction; they are an expression of its chemical behavior or conduct when undergoing chemical change. This method of expressing the classification is perhaps the better one. 12. Matter and Energy. A great deal can be learned from the simple experiments which were described earlier in the chapter. It will be recalled that the changes were brought about by means of heat, mechanical energy, light, and electricity. These are what are known as forms of energy. It is necessary to differ- entiate between matter and energy. A strictly accurate defini- tion of matter is difficult to formulate. Our experience and com- 12 INORGANIC CHEMISTRY FOR COLLEGES mon sense furnish us with a conception of matter. Matter occupies space, it has inertia, that is, it requires force to move it; it is the stuff of which the universe is made. Energy, on the other hand, is non-material; we become conscious of it only when it is associated with matter. A stone held in the air is different from the same stone resting on the earth; for by allowing the former to drop we can obtain work from it, drive a nail, or crush grain. The stone held away from the earth is said to have potential energy; it gives it up when it falls; and to raise it from the earth back to its original position, work must be done upon it. Energy manifests itself in work. In our experiments we have made use of four important forms of energy: heat, light, electricity, and mechanical energy. These can all be made to do work. The various forms of energy can be transformed, one into the other. For example, the mechanical energy of a water-fall can be transformed into electricity by means of a water-wheel and a dynamo. This can, in turn, be converted into light, heat, or mechanical energy. 13. Matter and energy are always associated. When any change occurs there is always a change in the energy; there may or may not be a change in the matter. From this point of view we can define physical and chemical change; if the change con- sists solely in energy it is physical, if the matter changes it is chemi- cal. Physics is thus the science which has to do with the study of changes in energy, either in amount or kind. As energy becomes manifest only through its action on matter, the physicist studies the behavior of substances which do not alter hi composition when energy changes take place in them. Physicists have studied, for example, the effect produced by moving a wire before a magnet, and have, as a consequence, invented the dynamo. They have investigated the behavior of light when it passes through glass, and the telescope is the result. In chemical changes both matter and energy change. The older, descriptive chemistry busied itself with the matter involved only. In modern chemistry, the energy changes which occur simultaneously with the changes in matter have been the subject of much study. This branch of the science is aptly called physical chemistry. For convenience it has been divided into several PHYSICAL AND CHEMICAL CHANGES 13 divisions: thermochemistry treats of the changes in heat energy when matter changes; photochemistry has to do with the rela- tionship between light and matter; in electrochemistry is studied the effect of electricity in producing chemical changes, as well as the production of electricity through chemical means; and other branches of the science consider the effect produced by mechanical energy. The subject of chemistry is such a broad one that it has been found desirable to study it from many points of view. Inorganic chemistry has to do with the substances found in the so-called mineral kingdom. The chief interest in organic chemistry centers in the compounds formed as the result of life-processes, and in the substances prepared in the laboratory from these compounds. Physiological chemistry, as the name implies, is the chemistry of the processes studied in physiology. Many other branches of the science are highly developed, such as industrial, metallurgical, mineralogical chemistry, etc. EXERCISES 1. State as fully as possible the accidental and characteristic properties of the following: (a) a crystal of salt, (6) a copper wire, and (c) a block of wood. 2. Study the experiments described to illustrate physical and chemical change (sections 6-9), and state the accidental and characteristic properties of the substances used and those obtained in each case. Show why the con- clusions reached are justified. 3. State which of the following are physical changes and which are chemi- cal changes. Examine each case from the point of view of characteristic properties, (a) tearing a piece of paper into small bits, (6) the falling of a stone, (c) burning of wood, (d) digestion of food, (e) lighting an incandescent lamp, (/) collision of two railroad trains, (gr) change of cider into vinegar, (h) making toast from bread, (i) exploding fire-works, (j) decay of a flower. 4. Name a physical and chemical change which occurs, (a) when a motor boat with a gasoline engine is running, (6) when a bell is rung by an electric battery, (c) when a lamp burns. 5. Name three physical and three chemical changes not mentioned in this book and give reasons for your conclusion in each case. 6. Name as fully as possible the physical properties of, (a) starch, (6) salt, (c) gold, (d) lead, and the chemical properties of (e) paper, (/) copper, (g) iron. 7. How could the following changes in energy be brought about: (a) heat 14 INORGANIC CHEMISTRY FOR COLLEGES into light, (6) heat into mechanical energy, (c) electricity into light, (d) mechanical energy into heat, (e) electricity into heat? 8. Name one change in each case not mentioned in the text which can be produced by (a) heat, (6) mechanical energy, (c) light, (d) electricity. State which of the changes are physical and which are chemical. 9. State what changes in energy take place in each of the following: (a) driving a nail with a hammer, (6) running an electric car uphill, (c) heating an electric flat iron. CHAPTER III ELEMENTS AND COMPOUNDS 14. As the result of the study of the last chapter the student has begun to learn how to observe and analyze phenomena. A closer inspection and a deeper understanding of the nature of the substances involved are necessary. Again, we shall seek the help of experiments in order to make clear the distinctions to be drawn. We shall repeatedly do this, for experimentation is the foundation upon which chemistry is based. We solve our problems, and clarify our conceptions by direct questions to nature; we bring together the things whose behavior we wish to understand, observe for ourselves the result, and draw our own conclusions. This is the so-called scientific method; it is a more or less novel one to the student whose point of view has been formed largely by reliance on authority. He has drawn his knowledge from books, and has had little opportunity to find out for himself. In chemistry he will learn to see with his own eyes; to answer many of his questions by obtaining first-hand knowledge in the laboratory and through the experimental demonstrations in the lecture-room. It should be constantly borne in mind that chem- istry is correctly studied in this way, and the student will be taught to realize the great value of such a method in all problems he may meet inside or outside the classroom. 15. Pure Substances and Mixtures. When we examine closely a piece of marble and a piece of granite we observe a marked difference between them. The one appears to be uniform in properties, the other not. We can distinguish in the granite three distinct substances: one is white, another is pink, and the third is black and glistens. Our common sense tells us that granite is a mixture, whereas marble appears to be a uniform substance. If we were to powder the two specimens we could separate the three constituents of the granite; in the case of the 15 16 INORGANIC CHEMISTRY FOR COLLEGES marble, however, each piece would be made up of the same material as every other piece. Granite is a mixture, and marble is what is called in chemistry a pure substance. In a pure substance every part is like every other part; the substance is uniform in properties. This refers, of course, to what we have called char- acteristic properties. Two pieces of marble might possess differ- ent sizes and shapes, and yet they are both marble; their character- istic properties are identical. We have chosen a very simple example to bring out the mean- ing of the words mixture and pure substance. We made use of the characteristic property of color in this case to come to the conclusion that granite is a mixture. If we were asked to deter- mine whether a given material uniform in color, is a pure sub- stance or a mixture, what would we do to aid us in answering the question? We must, evidently, make use of a characteristic property, other than that of color. If we had a mixture made up of sand and sugar, both in the form of a very fine powder, we could readily show a lack of uniformity by using solubility in water as the characteristic property. The sugar would dissolve, and the sand would not; the material is a mixture, for different parts of it possess different characteristic properties. 16. Another experiment clearly illustrates the point under discussion. Let us mix some iron and some sulphur, both finely powdered, and grind them together. We obtain as the result, what appears to the casual observer to be a uniform product. Is it a mixture or a pure substance? In this case we can make use of one of a number of characteristic properties to answer the question. Iron is gray, and sulphur is yellow. With a magnifying glass we can observe gray and yellow particles; the product is a mixture. Again, iron is attracted by a magnet, and sulphur is not. If we bring a magnet near the product, a part only clings to it. A third method can be used. Sulphur dissolves in a liquid which is called carbon disulphide, while iron does not dissolve. If we shake the product with this liquid and then pour the mixture on a piece of paper which is folded and placed in a funnel, a clear liquid will run through the paper and the iron will not. When this liquid is allowed to evaporate, it will be found that sulphur is left behind. We have by this means separated our product into two distinct parts, each of which is different from the other. ELEMENTS AND COMPOUNDS 17 17. Chemical Reactions. By carrying our experiment with iron and sulphur further, we can learn much more. A certain weight of sulphur and, of iron, which experience has shown are the correct amounts, are mixed and heated. The tube contain- ing the mixture is held in such a position that the bottom end of it, only, is placed in the flame of the burner. In a short time the mixture will get red hot where it is being heated. The tube is then removed from the flame. A striking phenomenon takes place; the brilliant glow slowly travels up the tube, and, finally, the whole mixture gives off light. When the tube is cold it is broken and the product is examined. Is it still a mixture of iron and sulphur? We can apply the tests used before. The prod- uct now is no longer a powder. In the lump which has been formed we cannot distinguish the fine particles of sulphur; a magnet has no effect on it; carbon disulphide does not dissolve from it yellow sulphur. A profound change has occurred; the product lacks the characteristic properties of iron, and of sulphur. The change has been a chemical one there has been a change in substance, and energy has been transformed; for as a result of the change light and heat were produced. A chemical reaction has occurred. What took place is called a reaction because the iron acted on the sulphur, and the sulphur acted on the iron. 18. Chemical Compounds. We next ask what has become of the iron and of the sulphur. The two substances have dis- appeared as such, and a new substance has been formed. The transformation has been most remarkable, for we can no longer recognize iron or sulphur in the product of the reaction. We explain this by saying the two substances have united chemically; a chemical compound has been formed. There was a marked change in the energy as well as in the matter when the reaction took place. In order to bring about the change quickly it was neces- sary only to start the reaction; it then proceeded of itself, and light and a large amount of heat were produced. Whence came this energy? Experience, based on more than a century of care- ful experimentation, teaches us that we cannot make energy; the amount in the universe is constant, and man can only change it from one form to another. As this is a fundamental truth, the energy transformed must have been associated with the iron and the sulphur. Energy in this form can be recognized only 18 INORGANIC CHEMISTRY FOR COLLEGES when the substances containing it undergo chemical change; it is then transformed, in part, into other forms of energy. We are thus led to add a new form of energy to those about which we have studied; it is called chemical energy. In the light of these conceptions we can describe more fully what occurs when iron reacts with sulphur. Iron as we know it consists of a certain form of matter with which is associated chemical energy; likewise sulphur, which possesses the properties we recognize, consists of a different form of matter associated with energy. When the two unite chemically a part of the energy is lost; it is transformed into light and heat. The resulting compound contains the material substance of the iron and that of the sulphur, but less energy than the two possessed before the change. The properties of substances change when we change the energy they contain, although the amount of material of which the substance is made does not alter. The compound formed when iron reacts with sulphur is called iron sulphide; the name is well chosen as it tells us that iron and sulphur are present in it. This compound is a pure substance; we can study it in many ways and we shall find that every part of it has the same characteristic properties possessed by every other part. When iron and sulphur are brought together we have seen that a mixture is formed; when this is heated a chemical reaction takes place, chemical energy is lost, being transformed into light and heat, and a chemical compound results. 19. One more interesting question remains to be answered. Can we separate the compound into its constituents by any means? Can we get back the iron and the sulphur? We shall see later that this can be done, but it is of importance to note that in order to do this we must make use of chemical reactions; we must give back to the iron and to the sulphur the energy which they lost when they united with each other. If we restore to the matter which is present in iron the right amount of chemical energy, it becomes metallic iron again. We thus see that it is possible to separate both mixtures and compounds into their constituents. In the case of mixtures we use mechanical or physical means; in the case of compounds we must resort to chemical changes. For this reason the two classes have been called mechanical mixtures ELEMENTS AND COMPOUNDS 19 and chemical compounds. We are again impressed with the necessity of a clear understanding of what is meant by a physical change and a chemical change. The discussion has been gone into at such length on account of its fundamental importance. 20. Elements. There is still more we can learn from the experiment with iron and sulphur. The substance formed as the result of their interaction contains both sulphur and iron; two distinct forms of matter enter into its composition. But what of the iron itself? Does it contain more than one substance? For many years chemists have studied iron and they have never obtained from it alone, more than one kind of matter. It is, therefore, called an element. Sulphur is also an element; it has never been broken up into two or more substances. Iron sulphide, on the other hand, is a compound, for by using the right methods it can be decomposed into iron and sulphur, the elements of which it is made up. A chemical compound has the characteristics of a simple substance, that is, it is uniform in properties; in addition, it contains two or more elements in chemical combination. It is important to note that the meaning of the word compound as used in chemistry, is quite different from that commonly assigned to it to the ordinary definition as given in the dictionary. Centuries of study have led to the conclusion that the universe is made up of about ninety elements. The infinite number of substances in the world have been produced as the result of the combination in different proportions of these elements. Only a small number of the elements are present in the things with which we come in contact. Less than a dozen are contained in the great variety of living matter around us. EXERCISES 1. State which of the following are pure substances and which mixtures: (a) salt, (6) vinegar, (c) sand, (d) sugar, (e) paper, (/) baking soda, (g) bread, (A) an egg. 2. How could you show that the following were mixtures and how would you separate each into its constituents? (a) chalk and salt, (6) powdered iron and powdered coal, (c) a shoe blacking made of soot and grease. 3. Name three substances that are elements and three that are compounds. 4. Write out definitions of the following using your own words : (a) physi- 20 INORGANIC CHEMISTRY FOR COLLEGES cal change, (6) chemical change, (c) mixture, (d} element, (e) chemical com- pound, (/) chemical energy. 5. State what changes in energy take place when (a) wood burns, (6) water falls, (c) electricity is made by a dynamo, (d) electricity is made by a dry cell, (e) an electric car moves, (/) toast is made on an electric stove, (g) magnesium is used in a flash-powder in photography at night, and (A) when gasoline is exploded in a motor. CHAPTER IV OXYGEN 21. Fire has been a source of fear, wonder, and veneration from prehistoric times; it has aroused the keenest interest and curiosity in man, and has been the subject of speculation by phi- losophers. It has received the closest study by investigators, and, as a result, its secret has been found out, and we have learned how to master it. The Greeks were of the opinion that the uni- verse was made up of four principles: earth, air, fire, and water. But the philosophy of the ancients did not get them very far toward the solution of the secrets of nature. They wasted time trying to solve the mysteries of natural processes by pure reason- ing alone. They disputed at great length, for example, as to whether matter was infinitely divisible, a purely philosophical question, but did not try to find out more about matter itself by examining it. One philosopher is said to have had his eyes put out, so that he might reason the more keenly and not be distracted by the crude material world. Such methods were futile, and after centuries of ignorance man learned the art of experimentation. Then the truth was discovered, the mystery disappeared, and fire ceased to be worshiped; it became the servant of man. If we are to produce fire by burning wood, we know air must be present. It was the discovery and isolation of the substance in the air which made burning possible, that solved the mystery of fire. Historical research has disclosed the fact that a number of chemists discovered at about the same time this important constituent of the air the substance which was later called oxygen. As early as 1489 the observation was made that when the sub- stance we now call the oxide of mercury is heated a " spirit " is given off; but to Priestley, an English scientist, is given the honor of making the discovery. It was he who published the fact to the world, and his observations were the basis of a correct explana- 21 22 INORGANIC CHEMISTRY FOR COLLEGES tion of combustion or burning. As long as a discovery is unknown to the world and is not used, it is of little value. For this reason the fact that Scheele, a Swedish chemist, recognized oxygen as a definite substance before its discovery by Priestley, is of historical interest only in connection with an account of Scheele's life and work. Priestley was a clergyman who had a most unusual curiosity in regard to nature; he took up natural sciences as a hobby, and made many discoveries of the greatest importance in chemistry. His theological speculations, which were recorded in many volumes, have been forgotten, but the results of his study of natural phe- nomena, which he carried out in the spirit of the amateur, are among the foundation stones of the science. Priestley happened to have a large burning lens, and amused himself by discover- ing the effect produced when, by means of it, the sun's rays were focused on various ob- jects. In one of his experi- ments he had some of the red oxide of mercury floating on the surface of metallic mer- cury contained in a glass tube (see Fig. 3) . When the heat was concentrated on the oxide of mercury it slowly FIG. 3. disappeared; the level of the mercury fell in thetube, which was found to contain an air, as Priestley said, or a gas as we say to- day. This gas was studied by him; substances burned in it with great brilliancy, and after it was found that a mouse breathed the new gas with perfect contentment, Priestley tried it on himself. He declared he was refreshed and exhilarated, but deprecated the use of the substance for this purpose, and said he believed we should be content with the air which nature furnished us. The discovery of this gas, which was made on August 1, 1774, OXYGEN 23 was of the greatest importance. In the hands of Lavoisier, a French chemist, it put chemistry on a true scientific basis. Lavoisier showed that this gas was present in the air, and that when many metals rusted or were heated in the air, they united with it. He showed, further, that if certain of the products so formed were heated, they decomposed into the metal and the gas. Lavoisier burned a number of substances in the gas, and found that when some of the resulting compounds were dissolved in water, acids were produced. For this reason he called the gas oxygen, deriving the name from two Greek words signifying acid- former. It is not only in combustion that oxygen plays such an impor- tant part. The element is essential to life; we breathe the air and the oxygen is carried by means of the lungs and blood into every part of our body, where it unites with the materials of the tissues. The chemical action which takes place produces heat, and is so regulated that the temperature of the body remains approximately constant under varying outside conditions. In the rusting of metals, oxygen takes part; it brings about decay in dead organic matter; it purifies the water in running streams; and it is present as a constituent of all living things and in many mineral sub- stances. Oxygen forms about 50 per cent of the atmosphere and crust of the earth; eight-ninths of water is oxygen. It is evident that it will be of the greatest interest to study oxygen in considerable detail. We shall learn how to obtain it in pure condition, study its properties and chemical behavior, and see how it can be put to a number of valuable uses. 22. Separation of Oxygen from the Air. The air, we shall see later, is a mixture of gases, one-fifth of any volume of it being oxygen. It is reasonable, therefore, to assume that air is a con- venient source from which to get oxygen directly. For a long time the gas could not be obtained in this way since there was no simple method known to separate gases from one another. When a method was devised to convert air into a liquid, the separation was conveniently effected. A liquid boils at a definite tempera- ture. For example, water is converted into a vapor or gas at 100 when the temperature is measured on a centigrade ther- mometer (81); wood alcohol boils at 66. When a mixture of the alcohol and water is heated to boiling, the alcohol changes 24 INORGANIC CHEMISTRY FOR COLLEGES into a gas more rapidly than the water; if we cool the vapor which comes off first to the temperature at which it liquefies, we obtain mostly alcohol. After the alcohol has boiled away the liquid left is pure water. When highly compressed air is cooled to a low temperature and allowed to expand rapidly, it changes into a liquid. The chief constituents of air are oxygen and nitrogen. The separation of the two substances is effected by boiling as in the case of wood alcohol and water. Nitrogen boils at a lower temperature than oxygen, consequently, when liquid air boils the nitrogen and a part of the oxygen pass off first. After the nitrogen has been converted into a gas, what is left, the residue, is oxygen. The method has become a practical one and furnishes both oxygen and nitrogen for commercial use. It requires elab- orate apparatus, and cannot be used by students in the laboratory. 23. Preparation of Oxygen by Heating Compounds Contain- ing It. (a) Mercuric Oxide. Oxygen is conveniently prepared by heating certain substances containing it. The use of mercuric oxide for this purpose has been seen to be of historical interest. The preparation can be conveniently carried out as follows: Some of the oxide is placed in a tube of hard glass, which does not soften when heated in a flame (Fig. 4). The tube is supported on a stand and connected by means of a piece of rubber tubing with a so-called delivery tube bent in such a way that any gas passing through it can be collected in an inverted bottle filled with water and placed over the end of the tube. As the gas comes from the delivery tube it forms bubbles in the water, which rise into the bottle and displace the liquid. The apparatus by means of which the gas is collected is called a pneumatic trough; the value of its use in collecting gases was first emphasized by Priestley. 1 When mercuric oxide is heated it decomposes into the ele- ments of which it is composed oxygen and mercury. The oxygen being a gas escapes and is collected in the bottle over water. As the heat is applied the mercury also passes into a gas, but changes into a liquid when it gets out of the region of the 1 Some of Priestley's original apparatus can now be seen in the National Museum in Washington. He was driven from England as the result of religous dissension, and settled in Northumberland, Pennsylvania, where he founded a college. OXYGEN 25 flame; it settles on the colder parts of the tube in drops, which can be recognized by their silver-like color. If the decomposi- tion is continued to the end, the red oxide disappears entirely. The gas collected can be shown to be oxygen by inserting into it a splinter of wood, the end of which is glowing. When a piece of wood containing sap is lighted, it burns with a flame. If this is extinguished by blowing on it or by moving it rapidly through the air, the wood continues to glow for some time; the end has the appearance of a hot coal. A characteristic property of oxygen is that it causes rapid combustion, and when a glowing splinter is inserted into it, a brilliant flame is produced as the wood burns rapidly. This effect is commonly used in testing for oxygen. 24. (6) From Potassium Chlorate. Mercuric oxide is an expen- sive substance, and cannot be used, therefore, to make oxygen on a large scale. Potassium chlorate is comparatively cheap. For a long time it was the principal source of the oxygen which was manufactured for commercial use, and it is still used for this purpose. It yields oxygen when heated, and is the material used in the laboratory in the preparation and study of the gas. The apparatus employed is like that just described. Potassium chlorate is familiarly called chlorate of potash; it is used as a wash in the case of sore- throat, and in certain tooth-pastes. When the sub- stance is heated it first melts; as the temperature is increased, it begins to decompose, and bubbles of oxygen rise through the melted mass; at a higher temperature the decomposition is rapid. 26 INORGANIC CHEMISTRY FOR COLLEGES When all the oxygen has been set free, the substance left is a white solid, which is called potassium chloride. Potassium chlorate contains three elements in chemical combination potassium, chlorine, and oxygen. The chemical change which takes place when it is heated to a high temperature consists solely in the libera- tion of the oxygen present. The resulting compound, potassium chloride, contains potassium and chlorine. 25. Catalytic Action. When certain substances are mixed with potassium chlorate, the rate at which the latter decomposes is increased. This can be shown to be true by a simple experiment. A test-tube is about one-fourth filled with potassium chlorate, and supported by means of an iron stand and clamp. The tube is next heated to the temperature at which decomposition takes place rapidly. It is then allowed to cool until the evolution of oxygen appears to have stopped and a glowing splinter of wood no longer bursts into a flame when introduced into the gas. If, now, a small quantity of powdered manganese dioxide the amount that can be held on the end of a blade of a pocket-knife is dropped into the tube, oxygen is rapidly given off. This can be shown by again testing with the glowing splinter; it will burn brilliantly. We can heat manganese dioxide, the substance that produced this striking result, up to the temperature used in the experiment and it yields no oxygen. Further, we can weigh carefully some manganese dioxide, use it to cause the rapid evolu- tion of oxygen from potassium chlorate, recover it and then weigh it again. The weight remains unchanged, and the manganese dioxide is unaltered. Many cases like that just described have been studied. A variety of substances are known which bring about reactions that do not appear to take place in their absence, although the substances themselves can be recovered unchanged after the reaction has taken place. Such substances are called catalytic agents or catalyzers, and the phenomenon is known as catalysis or catalytic action. It is probable that the part played by the catalyzer is to hasten a reaction which is already taking place very slowly. If a substance decreases the rate of a reaction it is called a negative catalyzer. 26. Many natural processes are brought about through catalytic action. The energy of the sunlight is stored up in a growing plant as a result of chemical actions induced by chlorophyl, OXYGEN 27 the green coloring matter in leaves. The digestion of food is accomplished as a result of the presence of catalytic agents in the saliva, the gastric juice, and other fluids of the body. Catalysis has been studied recently in great detail, and the process has been utilized in the preparation on a large scale of many substances of industrial importance. We shall have occasion to refer to a number of applications later, and when more facts are at command the \vay in which the phenomenon is thought to take place will be discussed. 27. Preparation of Oxygen by the Action of Electricity on Water. We have seen that heat is an important agency in effect- ing chemical change. When this form of energy fails, the desired result may often be produced by means of electricity. Water is a chemical compound of oxygen and hydrogen. When it is heated to an exceedingly high temperature it decomposes, in part, into the elements of which it is composed; but as these are both gases they cannot be separated readily, and the method is not used to prepare oxygen. When electricity, however, is passed through water it is broken down into hydrogen and oxygen without the application of heat. The process by which a substance is decom- posed by means of an electric current is called electrolysis. In order to render the water a conductor of electricity, pure water itself offering great resistance to the passage of the current, a small amount of sulphuric acid is added to it. The acid acts as a catalytic agent, for it can be recovered unchanged after a large amount of water has been decomposed. The experiment can be readily carried out in a Hofmann apparatus, which has already been described (9). When the decomposition has proceeded for some time it will be observed that the volume of the gas which has collected in one tube is exactly twice that in the other. If the gases are drawn off from the two tubes and tested separately, that which was formed in the smaller amount will be found to be oxygen. It will cause a glowing splinter to burst into flame. The other gas, which will be studied in detail later, is hydrogen; it will extinguish a lighted taper thrust into it, but will itself burn with an almost colorless flame. 28. When electricity passes through anything we say that the place from which the current flows is positive and the place to which it flows is negative; this is done as a matter of convenience 28 INORGANIC CHEMISTRY FOR COLLEGES in describing the phenomena produced by the current. In the Hofmann apparatus, for example, the current flows from the electrode in one tube, through the water, to the electrode in the other tube. The current passes in the solution from the pole at which oxygen is liberated to that at which hydrogen is set free. The electrode at which oxygen is set free is called the anode; hydrogen is liberated at the cathode. If we disconnect the wires where they are joined to the apparatus and reverse their positions, the poles at which the gases are evolved will be reversed. It is evident from this that the electricity which decomposed the water traveled in a definite direction; it is what is called a direct current of electricity. Oxygen and hydrogen are produced on the large scale for industrial purposes as the result of the decomposition of water by electricity. 29. Other Methods of Preparing Oxygen. Oxygen can be prepared from a number of compounds containing it. A sub- stance which is expensive, but is sometimes used in the laboratory on account of its convenience, is sodium peroxide. There are two compounds known which are made up of sodium and oxygen only. They differ in that one contains a greater proportion of oxygen than the other; the one containing the smaller amount of oxygen is called sodium oxide; the other is sodium peroxide. The latter, which is a white powder, is formed by burning metallic sodium in the air. When sodium peroxide is put into water, a reaction takes place. The products formed are sodium hydroxide, which dissolves in the water, and oxygen, which escapes. The presence of the hydroxide in the solution can be shown by the fact that the latter will change the color of red litmus paper to blue. Sodium peroxide which has been melted and poured into cans while in the liquid condition and allowed to solidify, is sold under the commercial name "oxone." In this form it is used in a specially constructed generator as a convenient source of small quantities of oxygen. 30. A number of substances other than those already men- tioned yield oxygen when heated to sufficiently high temperatures ; among these are the oxides of certain metals. The temperatures to which these compounds must be heated in order to bring about their decomposition into oxygen and the metal, vary widely. OXYGEN 29 The oxides of silver and gold yield oxygen at comparatively low temperatures. These metals do not react readily with oxygen; they do not rust; they are said to be chemically inactive. The oxide of magnesium, on the other hand, cannot be decomposed by heat alone at the highest temperature obtainable on the earth. We have seen that magnesium is an active metal; it burns rapidly in the air, and in doing so gives off a large amount of energy as heat and light. Magnesium oxide is a very stable substance on account of the fact that the elements of which it is composed lost so much of their chemical energy in their union. Substances which contain a large amount of chemical energy that can be transformed, are active, and enter readily into chemical change. Those which are formed as the result of evolution of a large amount of energy are more stable. Certain oxides lose a part of their oxygen when heated; among these are copper oxide and lead dioxide. Other substances which contain a large proportion of oxygen yield the element at a high temperature; but a consideration of these is deferred until they are described individually, since they are not commonly used as a source of the gas. 31. Properties of Oxygen. Oxygen is a tasteless, odorless, invisible gas. When subjected to a high pressure and a low temp- erature it changes to a bluish liquid, and at a still lower tempera- ture it becomes solid. Oxygen boils at 182.5 , 1 and melts at 227; the liquid or solid is attracted by a magnet. One liter of oxygen gas when weighed under standard conditions weighs 1.429 grams. Since the volume of a gas changes when the pres- sure on it changes, so-called standard conditions have been defined in order to facilitate comparisons between different gases; the standard temperature is and the standard pressure is that exerted by a column of mercury 760 mm. high. This pressure is the average pressure of the atmosphere at the level of the sea. Since one liter of air under these conditions weighs 1.293 grams, oxygen is slightly heavier than air. Oxygen dissolves sparingly in water; 100 volumes of the latter dissolve 4 volumes of the gas atO. 1 All temperatures given refer to the centigrade scale, which is described in Section 81. The metric system, which is used in expressing weights and volumes, is explained in the Appendix. 30 INORGANIC CHEMISTRY FOR COLLEGES 32. Chemical Conduct of Oxygen. The behavior of oxygen when it enters into chemical change can best be appreciated by considering some striking experiments. A number of large jars, filled with the gas, are provided. Into one is placed a glowing splinter; at once the wood burns rapidly with a brilliant flame. Into another is introduced a bit of charcoal; no action takes place. The charcoal is removed and heated to redness; when it is taken from the flame and allowed to remain in the air it loses its bright- ness, and quickly cools. It is heated to redness again and thrust into the oxygen; now it burns with a brilliant incandescence. The experiment is repeated with sulphur. When cold it does not react with the oxygen. When ignited it burns sluggishly in the air with a blue flame, but when introduced into oxygen a flame giving a bright blue light is produced. The burning of phosphorus is studied next. When the substance has been warmed it burns in oxygen with an intense white light. These experiments show that the substances which burn in the air, burn also in oxygen, but that the combustion takes place in this case much more rapidly; as a consequence the phenomena are most striking and brilliant. Some substances which do not ordinarily burn in the air do so when heated and placed in oxygen; iron furnishes a good example. The end of a file is heated to redness in a flame and put into a jar of oxygen; it burns rapidly, and sufficient heat is developed to melt the metal; and bits of it are thrown off in all directions like shooting stars. This behavior of iron is utilized in one form of fireworks. 33. A great deal can be learned from a careful study of these experiments. We observed that in order to bring about the reaction, in each case the substance introduced into the oxygen had to be previously heated. The temperature to which a sub- stance must be heated to bring about its rapid union with oxygen, with the evolution of light, is called its kindling tempera- ture. The word kindling as here used has the same significance as it has commonly. We kindle a fire by heating the materials to the temperature at which they continue to burn without the assistance of heat furnished from another source. The tempera- ture at which substances unite rapidly with oxygen is markedly affected by their physical condition; for example, a metal in the OXYGEN 31 state of a very fine powder has a kindling temperature much lower than the same metal when in the massive condition. Iron can be obtained in a form which takes fire spontaneously when brought into the air. We must next study the substances formed as the result of burning, or combustion. In the jars in which the wood and char- coal were burned we find nothing that we can see; they both contain an invisible gas. It is not oxygen, for if we introduce into it a glowing splinter it is at once extinguished. A new sub- stance has been formed. Both the wood and the charcoal con- tain the element carbon. The reaction which took place consisted in the chemical combination of this element with oxygen. The product is an oxide of carbon, and is called carbon dioxide. It is formed when any organic substance is burned, and is produced in our bodies, and is present in exhaled air. We are convinced that a new substance has been formed as the result of burning sulphur in oxygen by noting the odor of the gas that is left. If we breathe it, it chokes us and causes violent coughing; it will be recognized as the gas formed when sulphur burns in the air. We often suffer from it in improperly ventilated railroad stations, where the smoke from the locomotive is allowed to accumulate. The coal contains sulphur and when it is burned the gas formed from the latter escapes into the air. The sub- stance is formed as the result of the union of sulphur and oxygen; it is called sulphur dioxide. The jar in which the phosphorus was burned contains a white powder; it is phosphorus pentoxide. When the iron burned it was converted into a black solid, which is an oxide of iron. In all cases the chemical reaction which took place con- sisted in the direct union with oxygen of the substance burned; and an oxide was formed in each case. All compounds which con- tain oxygen and one other element are called oxides. The prefix used before the word oxide in naming compounds refers to the amount of oxygen present; this will be discussed in detail later. A careful study of the subject has shown that the products which are formed when carbon, sulphur, phosphorus, and iron are burned in oxygen, are the same as those produced when they are burned in the air. This fact proves to us that air contains oxygen. 32 INORGANIC CHEMISTRY FOR COLLEGES We can now understand as a result of our experiments what happens when a substance burns. Burning, or combustion as it is frequently called, consists in the union of a substance with oxygen with the simultaneous production of light and heat. The most striking chemical property of oxygen is shown in the act of combustion; we commonly state this by saying that oxygen sup- ports combustion. The substance which unites with oxygen is said to be oxidized; the process is called oxidation. Oxygen is an element. No one has ever been able to separate from it two or more substances. 34. Slow Oxidation. Many substances which burn unite slowly with oxygen without the production of light. If the same product is formed as the result of burning and of slow oxidation, the same amount of heat is given off in the two cases. When the oxidation is slow the heat is evolved at such a rate that there is no appreciable rise in temperature; it escapes as rapidly as it is produced and the substance does not get hot enough to emit light. An instruc- tive experiment will help to make this clear. Some steel wool l is placed in a long tube, which is closed at one end. The tube is supported in an upright position by a clamp, the open end being placed under water contained in a beaker. (See Fig. 5.) The tube con- tains air and iron. The level of the water in the tube is noted carefully, and the apparatus is not disturbed for some time. At the end of a day or two it will be seen that the iron has rusted, and that the water has risen in the tube. Iron oxide has been formed, and as the oxygen united with the iron and disappeared as a gas, the water rose to take its place. This experiment is instructive, as 1 Steel wool consists of long thin shreds of the metal matted together like wool. It is valuable for polishing wood, cleaning floors, etc. It is used in this experiment because it furnishes a large surface for the oxygen to act upon; the action takes place more rapidly as a result. FIG. 5. OXYGEN 33 it proves to us that when iron oxide is formed from iron and oxygen, both elements disappear as such. 35. Spontaneous Combustion. Certain substances unite so rapidly with oxygen that after a time enough heat is given off to raise their temperature to the kindling point, and light is produced as they burn. For example, if a piece of phosphorus is left in the air, it soon begins to give off a white smoke, which is phosphorus trioxide, produced as the result of the union of the element with oxygen. The chemical action takes place faster and faster; heat is produced rapidly, and, finally, the kindling point is reached and the phosphorus bursts into flame. Such a phenomenon as this is called spontaneous combustion. Certain substances which are not ordinarily spontaneously combustible exhibit this property under unusual circumstances. Cotton cloths that have been saturated with linseed oil, which is used in making paint, and left undisturbed for some time in a pile, have been known to take fire spontaneously. In this case the oil unites with oxygen from the atmosphere; as cotton is a poor con- ductor of heat, and the cloth inside the pile is protected from air currents which would carry away the heat produced as the result of the oxidation, the temperature slowly rises until a flame is produced. Cases have been reported of the spontaneous combustion of hay. The material was stored in a damp condition; decomposi- tion took place as the result of oxidation; and as hay is a poor conductor of heat, the latter accumulated in one place in suffi- cient amount to finally bring about active combustion. 36. Incombustible Substances. Compounds which are formed as the result of combustion do not, in general, burn. Substances like stones, bricks, asbestos, cement, and other so-called fire- proof materials, all contain large proportions of oxygen and cannot unite with more of the element; they are incombustible. 37. Uses of Oxygen. Great quantities of oxygen are con- sumed in the burning of coal and wood. In this case the source is, of course, the air, and we are apt to overlook the fact that oxygen plays such an important part in our happiness and comfort. We have no particular interest in the carbon dioxide formed when the coal burns, and let it escape into the air. What we desire is the energy produced; from it we obtain heat and power. Coal con- 34 INORGANIC CHEMISTRY FOR COLLEGES sists essentially of carbon. The element contains chemical energy. Oxygen contains chemical energy. When the two unite a part of the total energy of the two elements is transformed into heat. When we buy coal we buy chemical energy. When substances burn in oxygen a higher temperature is attained than when they burn in the air. Only about one-fifth of the air is oxygen; and when a substance burns in air a part of the heat produced is used to warm the substances present which do not play any part in the reaction; as a consequence the temp- erature does not rise so high as when pure oxygen is used. A number of uses are made of the high temperatures produced when substances burn in oxygen. Acetylene, the gas which was formerly much used as an illuminant on automobiles, produces a large amount of heat at a high temperature when it burns in oxygen. The form of apparatus used to produce a flame in this way is called an acetylene torch. By means of it it is possible to cut in a short time masses of steel that yield to the older methods only after the expenditure of much time and labor. The flame is slowly drawn across the metal; at the high temperature produced, the iron burns in the excess of oxygen, the oxide formed melts, flows away, and soon the metal is severed. In this way it is possible to cut rails and girders with dispatch. The mass of wreckage which resulted from the collapse of a great steel railroad bridge was cleared away with the aid of acetylene torches. By any other method the result would have been accomplished only at a great cost in time and labor. This is another interesting example of the use of the energy produced by chemical change. The form of carbon which is deposited in automobile cylinders does not burn in air, but is rapidly converted into carbon dioxide when ignited in an atmosphere of oxygen. Use is made of this fact in cleaning the cylinders of gasoline engines. Life is dependent on oxidation. The animal quickly dies if deprived of oxygen. In certain cases of extreme weakness in a patient, physicians resort to the use of oxygen. If breathing is enfeebled, it is advisable to inhale pure oxygen rather than air; four-fifths of the latter is an inert gas, which has no effect on the body. Oxygen is now employed with great success to restore persons rendered unconscious by noxious gases. Fire departments are OXYGEN 35 supplied with an apparatus, called a pulmotor, which is used to introduce oxygen into the lungs of persons who have been suffo- cated in a fire. 38. Oxygen in Nature. Having studied the chemical con- duct of oxygen we can now understand why it plays such an important part in nature. It is a very active element; it forms compounds with nearly all of the known elements. When the universe was being formed from the elements, most of them united with oxygen; and we find oxides of many elements in the earth's crust. Sand is made up chiefly of silicon dioxide; many ores of iron and other metals are oxides; and water, which is so abun- dantly distributed over the earth's surface, is an oxide of hydrogen. We also find a great variety of minerals which contain two or three elements in addition to oxygen; limestone, feldspar, clay, and many other substances which are present in the soil belong to this class. As has been noted, about one-half the earth's crust and the atmosphere around us is oxygen. Not only the earth, but the universe, as we know it, contains oxygen. Vast amounts of it are present in the sun a fact which is discovered by the study of the light which comes to the earth from this source. When all substances are heated to a sufficiently high temperature they give off light. Each element under these conditions emits light which is made up of certain colors the combination in each case being characteristic of the element. With the aid of a spe- cially devised optical instrument, called a spectroscope (615), the colors present in any light can be determined. The examination of the light given off by the sun shows that oxygen is present. In the decay of organic matter oxygen plays an important part. Through the agency of bacteria, which are organisms of, microscopic size, the carbon and other elements present are con- verted into oxides. The products of plant and of animal life are transformed in this manner, and find their way eventually into the atmosphere and the soil. A question naturally arises: If such great quantities of oxygen are consumed in combustion -and the growth and decay of living matter, is not the atmosphere losing it rapidly? And what will happen when this element, so necessary to life, has combined with other things? Nature has provided against such an emergency. When plants grow they absorb from the soil and the air the waste products of animal and 36 INORGANIC CHEMISTRY FOR COLLEGES vegetable life, and set free from these a large part of the oxygen, which in this way passes back into the atmosphere. In this wonderful transformation, which will be studied more fully later, the energy of the sunlight is changed into chemical energy which is stored up in the plant and the free oxygen. EXERCISES 1. State the effect on a furnace fire of (a) opening the door below the grate and (6) opening the door above the fire. Explain why the observed results occur. 2. What occurs when a wood fire is fanned or blown with a bellows, and when a burning match is placed in a strong draft of air or blown upon with the breath? Explain the reason for the different effects produced by the same cause. 3. When water is decomposed by electricity it is broken down into hydro- gen and oxygen. How do you think these gases could be changed into water again? 4. How could you tell which of the two poles of an electric battery to be used to decompose water, is the positive pole? 5. Sodium oxide and sodium peroxide are both solids. How could you tell them apart? 6. How could you distinguish potassium chlorate from potassium chloride? 7. Potassium chlorate is used in fireworks. Can you tell why? 8. In some flash-powders a mixture of aluminium and potassium chlorate is used. Explain the part played by each. 9. Powdered iron is used in a form of fireworks which produces a mass of brilliant sparks. What substances should be mixed with the iron to make it burn rapidly? 10. Of the oxides of the following metals which ones would you expect to yield oxygen when heated in a gas flame: iron, gold, platinum, silver, copper? 11. Gold and platinum are used to fill teeth. Why are these and not less expensive metals used? 12. If you were asked to examine jars each containing one of the following gases, how could you tell what substance was in each of the jars: oxygen, air, carbon dioxide, and sulphur dioxide? 13. If you were asked to prepare copper oxide and zinc oxide how would you do it? 14. If you were given a sample of silver oxide how could you prove it was not an element? 15. How could you prevent a large bulk of oily waste from taking fire spontaneously? 16. Why do we turn down the wick of a lamp that is smoking to stop the formation of soot? J.7. Why does not gas that is lighted burn back into the pipe? OXYGEN 37 18. If a bit of wood is held just inside the top of a lamp chimney it does not burn with a flame. Why? 19. Why will most plants not grow in the dark? 20. What is the original source of the energy produced by burning wood? 21. When iron burns in oxygen the elements lose chemical energy which is changed into heat. When the iron oxide is heated with carbon (charcoal) the metal is recovered and possesses its original amount of energy. Where did the latter come from and what evidence is there of the truth of your con- clusion? 22. Water contains 88.8 per cent oxygen. How many grams of oxygen can be obtained from (a) 100 grams, (6) 25 grams, and (c) 21.63 grams of water? How many liters of oxygen are obtained in each case? 23. Potassium chlorate contains 39. 16 per cent of oxygen, (a) How much oxygen is contained in 50.2 grams of potassium chlorate? (6) How much potassium chlorate must be taken to get 20 grams of oxygen? CHAPTER V HYDROGEN Hydrogen, which we have seen was formed along with oxygen as the result of the electrolysis of water, is an important element. It is a colorless gas, which is characterized by the fact that it is the lightest substance known. 39. Occurrence of Hydrogen. As hydrogen is a constituent of all living things it is widely distributed in nature. It occurs in the free condition that is, not in combination with any other element in the gases that issue from volcanoes and, in small quantities, in natural gas. Free hydrogen is an important con- stituent of the gas that is manufactured from coal and is used for the production of light and heat. Hydrogen is formed as the result of the decomposition of organic matter brought about through the action of certain bacteria. The occurrence of the hydrogen which is combined with other elements, is of greater importance from the chemist's point of view. It constitutes by weight one-ninth of water, and is present in all substances known as acids, and in many chemical com- pounds. It is a necessary constituent of living things. 40. Early History of Hydrogen. In the sixteenth century Paracelsus 1 noted the fact that when iron was treated with a dilute acid " an air rises which bursts forth like the wind." It was not until 1766, however, that the gas was really discovered and studied. In this year Cavendish, an English physicist, prepared hydrogen by the action of hydrochloric or sulphuric acid on zinc, tin, or 1 Paracelsus was a professor of medicine in the University of Basle, Switz- erland. He was the first to discard the alchemical doctrines and aims, and to put aside the methods based on magic and superstition, which were used at that time in treating disease. He studied the action of chemical com- pounds on the body and used them as drugs. His influence was great, and men turned from attempts to make the philosopher's stone to a more prac- tical study of chemical substances. 38 HYDROGEN 39 iron. He isolated the gas and described its properties. In 1781 he showed that when hydrogen burned water was formed. Some time later, Lavoisier named the gas hydrogen, deriving the name from the Greek words signifying water-former. It was previously called inflammable air. 41. Preparation of Hydrogen, (a) By the Electrolysis of Water. We have already seen how oxygen and hydrogen can be prepared by the electrolysis of water (27), and that the gases are formed in the ratio of one volume of the former to two of the latter. We also discovered that hydrogen burns with a flame which is almost invisible, and that it does not support combustion. If some air is allowed to mix with the gas before it is lighted, the reaction takes place with the production of a characteristic sound that is produced as the result of an explosion. This behavior of the gas burning when pare with an almost colorless flame, extinguishing a glowing splinter, and exploding when mixed with air and ignited is used as a convenient test for hydrogen. 42. (6) Preparation of Hydrogen by the Action of Sodium on Water. The methods which have been described up to this point to separate oxygen and hydrogen from their compounds, are based on the decomposition of the latter through the action of heat or electricity; energy was required to bring about the change. We are now about to see how chemical energy can be used for this purpose. When a piece of sodium is put upon the surface of water, a violent reaction takes place; hydrogen is rapidly evolved, and, as a consequence of the evolution of a large amount of heat, the metal melts, and finally catches on fire and burns. Sodium is a very active element, and contains a large amount of chemical energy, which exhibits itself when the metal is brought into contact with oxygen or certain compounds containing oxygen. In order to regulate the reaction between sodium and water a simple device can be used. A small piece of the metal is wrapped in copper gauze; when this is put into water the weight of the gauze causes it to sink; as a consequence, the sodium cannot come into contact with the oxygen of the air, and cannot burn. The gas evolved is collected by holding a tube filled with water over the stream of bubbles as they rise. It is shown to be hydro- gen by inserting into it a burning stick of wood; the gas burns, and the flame produced by the burning wood is extinguished. 40 INORGANIC CHEMISTRY FOR COLLEGES The chemical change which takes place in this striking reac- tion is this: The sodium unites with the oxygen present in the water, and drives out a part of the hydrogen. The compound of sodium, hydrogen, and oxygen which is formed is called sodium hydroxide, the name indicating the presence of the three elements. The substance is a white solid that dissolves in water. Many compounds which contain hydrogen and oxygen are called hydrox- ides; we have, for example, iron hydroxide, tin hydroxide, etc. The metallic hydroxides which dissolve in water, such as sodium hydroxide, show the characteristic property of changing the color of a dye called litmus from pink to blue. 43. (c) Preparation of Hydrogen by the Action of Certain Metals on Water. Elements other than sodium are able to take the oxygen from water and thus set free hydrogen. Whether or not any given element can do this is determined by the chemical energy the element possesses. Active elements like magnesium, iron, and zinc decompose water. In these cases, however, it is necessary to apply heat to bring about the reaction, or to hasten it so that the decomposition of the water takes place with reason- able rapidity. The less active the metal, the higher it must be heated before reaction occurs. Magnesium decomposes water slowly at its boiling-point, zinc at a higher temperature, and iron must be heated with steam to dull redness before hydrogen is freely evolved. An experiment to demonstrate the formation of hydrogen in this way can be carried out easily. A supply of steam is furnished by boiling water in a flask, which is connected with one end of an iron tube filled with iron filings; to the other end is joined a delivery tube which leads to a pneumatic trough (see page 42 for a drawing of the latter). The iron tube is supported in a furnace supplied with a number of gas burners, and is heated to a high temperature. As the water boils steam passes over the iron filings; hydrogen and an oxide of iron are formed, and the gas escapes through the delivery tube and is collected over water in the pneumatic trough. The most inactive metals like gold, silver, and platinum do not decompose water at any temperature, since their affinity for oxygen is very small compared with that of hydrogen for oxygen. It takes a large amount of energy to separate hydrogen from oxygen when these elements exist in chemical combination as HYDROGEN 41 water, and the inactive metals are not able to furnish this energy. 44. (d) Preparation of Hydrogen by the Action of Certain Metals on Acids. We shall have occasion to use acids repeatedly in the preparation and study of a number of substances, before these acids can be conveniently discussed in detail. For a correct understanding of the chemical reactions into which they enter, it is necessary to learn something in a general way about a few of these compounds. A characteristic property of acids with which we are all familiar is their sour taste. Vinegar contains an acid acetic acid which, when separated in a pure condition, is a color- less liquid possessing an odor which we recognize. Sulphuric acid is a heavy oily liquid, and is sometimes called oil of vitriol. A solution of the acid in water is commonly called sulphuric acid, and the substance unmixed with water, concentrated sulphuric acid. Sulphuric acid contains the elements hydrogen, oxygen, and sulphur. Hydrochloric acid is another reagent which is frequently used. It consists of a solution in water of hydrogen chloride, which is a compound of the elements hydrogen and chlorine. A saturated solution of hydrogen chloride is known as concentrated hydro- chloric acid; if more water is added to this we have dilute hydro- chloric acid. Hydrogen is present in all acids; and all acids yield hydrogen in the free condition when they are treated with certain metals. The gas is most conveniently prepared by the action of dilute hydrochloric acid on zinc. When a piece of the metal is covered with the acid, reaction soon begins, bubbles of gas appear and rise through the liquid, and the solution grows warm. As the tempera- ture rises the rate at which hydrogen is set free increases rapidly. 45. If we wish to collect the gas for study, the reaction is carried out in a hydrogen generator. This consists of a bottle, closed by a stopper, through which two tubes pass. One of these is what is called a safety tube, or thistle tube (Fig. 6) . The tube is so adjusted that one end of it nearly touches the bottom of the bottle. The second tube passes just through the stopper and is bent at such an angle that it can be conveniently connected with a delivery tube. Some zinc is placed in the bottle, the stopper is put in place, and acid is poured in through the safety tube, care 42 INORGANIC CHEMISTRY FOR COLLEGES being taken that the end of the latter is beneath the surface of the liquid. After reaction has proceeded for some time, and the air which was in the bottle has been driven out, the delivery tube is placed under the mouth of a bottle which has been filled with water and inverted in a pneumatic trough. If hydrogen stops coming off as a result of the exhaustion of the acid, more of the latter can be added through the safety tube without admitting air to the generator. This form of apparatus is simple in construction and is used in the preparation of a number of gases. FIG. 6. It has been stated that hydrochloric acid is a compound of hydrogen and chlorine. When zinc reacts with the acid, the hydrogen is liberated and the metal unites with the chlorine; the compound formed as a result is called zinc chloride. If the water which is present is boiled away the compound is left as a white solid. When zinc reacts with sulphuric acid, hydrogen is set free, and the metal unites with the oxygen and ;ulphur which were present in the acid; the compound formed in this case is called zinc sulphate; it, too, is a white solid, soluble in water. Zinc and acetic acid give zinc acetate and hydrogen. 46. A number of metals in addition to zinc liberate hydrogen from acids; among these are magnesium, aluminium, iron, and tin, all of which are more or less active metals. Copper, silver, HYDROGEN 43 and gold, on the other hand, do not possess this power. The rate at which the reaction takes place in the case of several metals varies with their activity, provided other conditions are the same. This can be readily shown by putting pieces of magnesium, zinc, and tin of equal size and shape into hydrochloric acid. The rate of evolution of hydrogen decreases with the metals in the order given. 47. The rate at which different samples of any one metal lib- erate hydrogen from an acid is markedly affected by the purity of the material. Zinc of the highest purity, for example, reacts with sulphuric acid with extreme slowness; impure zinc, which contains small amounts of carbon and other substances, dissolves readily in the acid. The impurities act evidently as catalytic agents. A simple experiment will show this effect. If a piece of pure zinc is put into dilute sulphuric acid there is scarcely any action. If now, a piece of platinum is placed in contact with the zinc, hydrogen is evolved. The platinum is recovered unchanged after the zinc has dissolved; it served as a catalytic agent. Other metals than platinum can be used for this purpose. The follow- ing is an instructive experiment: Into each of two test-tubes con- taining sulphuric acid is placed a piece of zinc which reacts slowly with the acid, the pieces of metal used being approximately of the same size and shape. To one tube is added a few drops of a solu- tion of copper sulphate. We watch carefully what happens: A black deposit appears on the zinc in the tube to which the copper sulphate was added. As soon as this is formed the rate at which hydrogen is evolved increases markedly. At the end of a short time the reaction which has been catalyzed takes place with con- siderable rapidity, whereas the other proceeds sluggishly, if at all. The black deposit formed when zinc is treated with a solution of copper sulphate is metallic copper; the color of the latter is due to the fact that it is in the form of an exceedingly fine powder. In this form it is very active as a catalytic agent on account of the fact that the surface of the metal is great, and there are very many points of contact between the copper and the zinc. A small amount of copper sulphate is often added to hydrogen generators to increase the rate at which the gas is formed. 48. (e) Preparation of Hydrogen by the Action of Certain Metals on Bases. We have learned something of sodium hydroxide 44 INORGANIC CHEMISTRY FOR COLLEGES in the study of the action of sodium on water. It will be recalled that hydrogen and this compound were formed. Sodium hydrox- ide belongs to a very important class of compounds called bases; Hydroxides, as the name implies, contain hydrogen and oxygen. From some of these hydrogen can be readily obtained by the action of the more active metals. We can trace again the cause of the action to the affinity of the metals for oxygen; they seek out this element, and displace the hydrogen united to it. When aluminium is put into a solution of sodium hydroxide, hydrogen is set free. The aluminium unites with the oxygen and sodium and forms a compound called sodium aluminate, which is soluble in water. When zinc is heated with solid sodium hydroxide a similar reaction takes place; hydrogen and sodium zincate are formed. 49. Physical Properties of Hydrogen. Hydrogen is a color- less, tasteless, and odorless gas. The odor observed when hydro- gen is prepared by means of impure metals, such as commercial iron, is the result of the presence of impurities in the gas. Hydro- gen has been liquefied and solidified; it boils at 252.5 and melts at 260. Hydrogen is the lightest substance known; one liter of it under standard conditions weighs 0.08987 gram. It is less than one-fourteenth as heavy as air. If we wish to pour hydrogen from one vessel into another, it is neces- sary to hold them in an inverted position (Fig. 7); as the bottom of one vessel (a) is lowered in the di- rection indicated by the arrow 6, air rushes into it and the hydrogen forced out ascends into the second vessel from which the air is expelled. If the gas in c is now tested it will be found to contain hydrogen. The lightness of the gas can also be shown by filling soap bubbles with it. When they are released they rise rapidly in the air, whereas bubbles filled in the ordinary way with air FIG. 7. HYDROGEN 45 from the lungs are heavier than air and sink to the ground. If the bubbles containing hydrogen are brought into contact with a flame as they rise into the air, they explode with a loud noise. Hydrogen is adsorbed by many metals, the amount of the gas taken up in any case being affected by the physical condition of the metal. Platinum in the form of a very fine powder adsorbs about fifty times its volume of the gas; palladium when properly prepared will adsorb over 500 volumes of hydrogen. 50. If a jar containing hydrogen is inverted and placed mouth to mouth over another jar containing air, and the two are allowed to stand undisturbed for a few minutes and are then examined, it will be found that the gases have mixed. This can be shown by inserting two glass plates between the jars and then separating them; when the gases are tested by means of a lighted taper, an explosion occurs in each case. Since hydrogen is so much lighter than air we might expect it to remain in the upper vessel to float on the air as a cork does in water. The fact that it does not is an illustration of an important property of all gases; when they are brought together they mingle and finally a uniform mixture results; this property is called diffusion. Graham, a Scotch chemist, who studied this phenomenon carefully, discovered a simple law which expresses the relative rates at which gases diffuse. The rates at which gases diffuse are inversely proportional to the square roots of their densities. The lighter the gas the faster it diffuses. One liter of hydrogen and of air weigh 0.08987 gram and 1.293 grams respectively. The rates at which they diffuse are as Vl.293 is to V0.08987 as 3.8 is to 1. Oxygen is sixteen times as heavy as hydrogen; it diffuses one-quarter as rapidly as hydrogen. 51. The fact that hydrogen diffuses more rapidly than air can be shown by a simple experiment. A cup of unglazed porcelain attached by means of a stopper to a glass tube, is inverted and the lower end of the latter placed in a beaker containing water (Fig. 8). A jar containing hydrogen is next placed over the porce- lain cup. The hydrogen diffuses into the cup through the pores of the latter and the air diffuses out. Air is forced out of the cup through the tube and rises through the water in bubbles; this takes place because the hydrogen enters the cup through the pores more rapidly than the air passes out in this way. The gas 46 INORGANIC CHEMISTRY FOR COLLEGES accumulates, as a consequence, and finally forces its way out through the water. 52. Chemical Conduct of Hydrogen. The most striking chemi- cal property of hydrogen has been exhibited in the experiments which have been described. The gas burns with a flame that can scarcely be seen. If any solid material is allowed to pass through the flame, it is heated to such a temperature that it gives off light. A convenient way to do this is to shake together near the flame two blackboard erasers which contain crayon dust; as the parti- cles of the latter enter the flame they become incandescent and give off yellow light. When hy- drogen burns it unites with, oxy- gen and forms an oxide. It will be shown later that this oxide is water. Hydrogen does not unite readily with most of the sub- stances which burn in oxygen. When a burning splinter is insert- ed into the gas it is extinguished, because wood does not unite with FIG. 8. hydrogen. The gas is said not to support combustion. When this experiment is carried out, the hydrogen itself burns at the mouth of the vessel containing it. When hydrogen and oxygen are mixed and kept at the ordi- nary temperature no apparent reaction takes place between them. If the temperature of the mixture is raised, reaction sets in very slowly at about 300, and at 700 the gases unite with explosive violence. The temperature at which the union of the two elements takes place is affected by the nature of the substance of which the vessel containing the elements is made. The rate of the reaction is markedly influenced by catalytic agents. An experiment illustrates this clearly. If some finely divided platinum at the temperature of the room is introduced into a stream of hydrogen HYDROGEN 47 issuing from a generator, the gas immediately bursts into a flame. Platinum for this purpose is prepared by soaking some asbestos, a non-inflammable mineral substance, in a solution of platinum chloride. The asbestos is then heated to a high temperature; the platinum chloride adhering to it is decomposed by the heat, the metal is deposited as a very fine powder or sponge, and the chlorine escapes into the air. Platinum prepared in this way is an active catalytic agent. It will be recalled that copper in a fine state of division was used to catalyze the reaction between zinc and sulphuric acid. As the catalytic action takes place at the surface of the metals, it is evident that the greater the surface the greater the catalytic effect. 53. When a mixture of hydrogen and oxygen in the right proportions is ignited, the union of the elements takes place with explosive violence. This fact has often led to accidents when attempts have been made to light the gas issuing from a hydrogen generator. When the acid is poured upon the metal and action begins, the apparatus is full of air. As the reaction proceeds the gas forced out of the generator is a mixture of hydrogen and air, the proportion of hydrogen increasing as the metal dissolves. If a flame is brought in contact with this mixture of gases, it will be ignited, the flame will travel back into the generator and a violent explosion will occur; and the flask will probably be shattered. An experiment to illustrate this can be carried out with safety by covering with a strong box the hydrogen generator as soon as the acid is introduced into it; through a hole in the side of the box is placed the tube through which the gases issue. A lighted burner is placed near the end of the tube. When the mixture of gases contains its constituents in the right proportions, it ignites; the explosion which results produces a loud noise and shatters the generator. The range of explosibility of mixtures of hydrogen and air is wide; if such a mixture contains from 9.5 to 65 per cent by volume of hydrogen it will explode when ignited. 54. An explosion is produced as the result of the rapid forma- tion of gases or their sudden expansion. When a mixture of oxy- gen and hydrogen is ignited a large amount of heat is generated in an incredibly short time. Gases increase in volume when they are heated. As a consequence, the water-vapor formed as the result of the union of the oxygen and the hydrogen expands 48 INORGANIC CHEMISTRY FOR COLLEGES rapidly. Since the heat is generated so quickly its full effect is produced, and there is not time for it to be lost to the surround- ings. If the explosion takes place in an open vessel with a wide mouth, the gases can expand into the air, and the vessel is not shattered. The air is set in rapid motion and a sound is produced as a result. If the vessel containing the exploding mixture is closed, it offers a resistance to the tendency of the gas to expand, and pressure is produced. Whether or not the containing vessel is destroyed depends upon whether it is strong enough to resist the pressure generated as a result of the explosion. The vapor produced from substances which burn, such as gasoline and ether, form explosive mixtures with air. When an explosion occurs the hydrogen in such compounds unites with oxygen to form water- vapor; if they contain carbon, carbon dioxide, which is a gas, is formed. Hydrogen unites directly with a number of other elements, such as chlorine, sulphur, sodium, etc.; the reactions will be con- sidered when these elements are discussed. 65. Hydrogen as a Reducing Agent. Hydrogen has a strong tendency to unite with oxygen. It will unite not only with free oxygen as we have seen, but also with oxygen which is held in chemical combination by other elements. An experiment will illustrate clearly this conduct of hydrogen. A hydrogen generator is attached to one end of a straight tube made of hard glass, which is so placed that it can be heated conveniently by a number of gas burners; to the other end is joined a glass tube in the shape of the letter U, which is kept cold by immersion in water. Some copper oxide is put in the straight tube, and hydrogen from the generator is passed through the apparatus. When the air has been expelled the copper oxide is heated. Reaction soon begins; the black copper oxide changes to metallic copper, which can be recognized by its characteristic color; and water collects in the U-tube. We see from this experiment that the products of the reaction between hydrogen and copper oxide are copper and water; it consists in the transfer of the oxygen from one element to the other. The copper oxide is said to have been reduced. Reduc- tion is the name applied to the process by which oxygen is removed from a compound in which it is present. The substance which effects the change is called a reducing agent. One of the most HYDROGEN 49 important chemical characteristics of hydrogen is its reducing power. In this chemical reaction the oxygen that is removed from the copper oxide unites with the hydrogen, which is oxidized as a result. In general, reduction and oxidation take place simul- taneously; one substance is reduced and the reducing agent is oxidized. The oxides of many elements can be reduced by hydrogen. The ease with which the reduction takes place is determined by the activity of the element present in the oxide. In the case of the most active metals the transfer of oxygen cannot be effected; for example, magnesium oxide and aluminium oxide cannot be reduced by hydrogen. Processes of oxidation and reduction play a very important part in chemistry; many metals are extracted from their ores by reduction. Iron is obtained, for example, by heating with carbon oxides of iron which occur in nature; the carbon serves as a reducing agent as it unites with the oxygen and thus sets free the iron. 56. Uses of Hydrogen. The uses of hydrogen are based upon its property of extreme lightness, upon the fact that it contains a large amount of chemical energy which is transformed into heat when the gas burns, and upon its reducing power. Most of the hy- drogen used commercially is obtained as a by-product in the manu- facture of caustic soda and chlorine by the electrolysis of solutions of sodium chloride. Hydrogen is used in balloons, but on account of the cost of its production it is frequently replaced by a gas manufactured by a special process from coal. The oxy-hydrogen blow-pipe is a form of apparatus designed to burn hydrogen in oxygen without the possibility of an explosion. The blow-pipe consists, in brief, of a tube through which oxygen passes, surrounded by a second tube that delivers hydrogen. The gases mix at the ends of the two tubes as they pass into the air. When the mixture is lighted an exceedingly hot flame is produced; its temperature is said to be 2500, whereas the temperature of a Bunsen flame produced by burning a mixture of coal gas and air is about 1500. The oxy-hydrogen flame is used when a source of heat at a high temperature, which can be conveniently manipu- lated, is desired. It is used to melt and work platinum and other refractory metals. It finds an important application in the 50 INORGANIC CHEMISTRY FOR COLLEGES making of apparatus from quartz or silica, a substance which resembles sand in chemical composition. Quartz cannot be melted in an ordinary gas flame. Apparatus made from it, such as tubes, dishes, retorts, flasks, etc., resists the highest heat commonly used in the chemical laboratory; on account of this and the fact that silica apparatus possesses other valuable properties, it is much used. 57. When substances which do not burn are heated to a high temperature they give off light, the brilliancy of which increases very rapidly with rise in temperature. This principle is made use of in the so-called calcium-light or lime-light. In an appropriate apparatus a piece of lime, which is calcium oxide, is heated by an oxy-hydrogen flame. Lime serves the purpose well, because at the temperature of the flame it does not melt; it emits an intense white light, which was formerly much used in projecting stereopti- con pictures on a screen. On account of the greater ease of hand- ling electricity, it has largely replaced the calcium light for this purpose, although the steadiness and other qualities of the latter still recommend it. Ordinarily illuminating gas, which contains a large amount of free hydrogen, is used in place of the pure gas, on account of the ease with which it can be obtained. The gases used in connection with the lime-light are stored in iron cylinders under pressure; the transportation of these is a source of inconvenience. Hydrogen is one of the most important constituents of illumi- nating gas, which is used as a source of light, heat, and power. The hydrogen is obtained in one form of this gas by passing steam over hot coal; the carbon of the coal removes the oxygen from the water, and sets hydrogen free. The reaction in this case is similar to that of the action of steam on iron, which has been studied. 58. Hydrogen is much used as a reducing agent in commercial chemistry. Many substances which are transformed into dyes are reduced by hydrogen during the process. Large quantities of hydrogen are used in making solid cooking fats from liquid vegetable oils, such as cotton-seed oil. When these oils are treated under pressure with hydrogen in the presence of a catalytic agent, the oil and the hydrogen unite and form a substance which is solid. Finely divided nickel is the catalytic agent used; this metal has been found to be very valuable to catalyze reactions in HYDROGEN 51 which hydrogen is involved. The direct addition of hydrogen to other substances is called hydrogenation. Large quantities of hydrogen are used in the manufacture of ammonia. In this process, which was developed in Germany during the recent war, hydrogen and nitrogen under pressure are brought into reaction through the agency of a catalyst. From the ammonia so prepared nitric acid was manufactured for use in the preparation of explosives. EXERCISES 1. Would you expect (a) water to be decomposed when it is heated with mercury and (6) mercury oxide to be reduced to the metal when heated with hydrogen? 2. How could you show that sour milk contained an acid? 3. Name the compounds formed when iron, zinc, tin, lead, and aluminium are dissolved in (a) hydrochloric acid, and (6) in sulphuric acid. 4. What is the source of the heat produced when a metal dissolves in an acid? 5. How could you distinguish by chemical means (a) zinc from silver, (6) magnesium from platinum, (c) tin from silver, (d) aluminium from zinc? 6. How could you show that paper contains the element hydrogen? 7. What is the source of the hydrogen that is present in combination in a piece of wood? 8. (a) Why would you expect a mixture of air and gasoline vapor to explode? (6) What use is made of this reaction? (c) Why should gasoline be stored in a place where there is a free access of air? 9. In testing samples of hydrogen from a generator to find out if air is present, the tube containing the gas is held with the mouth down until it has been ignited. Why? 10. Give examples of common phenomena which are produced as the result of the fact that gases diffuse. 11. If a porous substance like silk is used as the fabric in making balloons to be filled with hydrogen, the gas in the balloon slowly becomes heavier and its lifting power is reduced. Why? This result is largely avoided by oiling the silk. Why? 12. One liter of air weighs 1.29 grams and 1 liter of hydrogen 0.09 gram. A balloon was constructed which held 1000 cubic meters of hydrogen, and weighed 600 kilograms when inflated. What weight would the balloon just lift from the eprth, assuming that it displaced 1000 cubic meters of air? CHAPTER VI THE ATOMIC THEORY. CHEMICAL EQUATIONS 59. Chemistry advanced very rapidly as soon as the balance was used as an aid in the study of natural phenomena. Before this time fanciful theories had been put forward in attempts to explain burning and rusting; combustible substances and metals that rusted were supposed to contain the spirit of fire, called phlogiston, which escaped when these remarkable changes took place. When it was shown, however, that a metal on rusting increased in weight, the view that this change was the result of the escape of phlogiston was no longer tenable. Lavoisier, who studied phenomena of this kind, showed not only that when a metal rusted it increased in weight, but that this increase was equal to the weight of the oxygen which united with the metal; the discovery of oxygen and the use of the balance thus fur- nished the true explanation. Lavoisier's work impressed upon scientists the necessity of quantitative measurements in the study of the changes that take place in matter; it resulted further, as we shall soon see, in the discovery of one of the most fundamental laws of nature the law of the conservation of matter. We have learned that when iron burns an oxide of iron is formed, and that when iron rusts an oxide of the metal is also formed. A close examination of the oxides in the two cases brings out the fact that they are different substances. How can this be? A quantitative study of the two reactions that is, a determina- tion of the quantities of the iron and the oxygen that react in the two cases will answer the question. The two compounds differ as the result of the fact that they contain iron and oxygen in differ- ent proportions the question is answered by an appeal to the balance. There are two oxides of hydrogen; quantitative analy- sis shows us that water contains a smaller percentage of oxygen than the second oxide, hydrogen peroxide. A study of quanti- 52 THE ATOMIC THEORY. CHEMICAL EQUATIONS 53 tative relations not only makes clear the composition of substances, but serves many practical ends. How much zinc must react with hydrochloric acid to produce six jars of hydrogen to be used for experimental purposes? How much cream of tartar should be mixed with a certain weight of cooking soda to make baking- powder? Such questions as these can be answered as a result of the study of the quantities of substances involved in chemical reactions. Up to this point we have emphasized the qualitative side the side which has to do with the different kinds of matter that undergo change. We shall now see that our knowledge will be greatly increased by a study of the quantitative aspect of these transformations. 60. Law of the Conservation of Matter. Lavoisier and others studied many chemical changes quantitatively; as a result, it was found that the sum of the weights of the substances entering into reaction equaled exactly the sum of the weights of the products of the change no matter was lost or gained. This conclusion has been repeatedly confirmed. Some years ago the German chemist Landolt carried out a long series of experiments to test the accu- racy of this conclusion; he worked with the greatest skill and used the most delicate instruments that could be devised for de- tecting changes in weight. Landolt placed the substances that were to interact in a tube shaped like an inverted letter V; one sub- stance was placed in one arm of the tube, the second substance in the other. The tube was next sealed to prevent any loss, and was carefully weighed. It was then inverted; the substances contained in it mixed, and the chemical reaction took place. After some time, when the tube returned to its original tempera- ture and the external conditions were the same as before the reaction took place, the tube was weighed again. No change in weight was observed. It seemed of importance to test with the greatest attainable accuracy the conclusion that there is no change in the total weight of the substances which are undergoing chemi- cal change, for this conclusion is the basis for all quantitative study in chemistry. This fundamental fact is summed up in the law of the conservation of matter, which is stated in various ways. One statement is that the total amount of matter in the uni- verse does not change; another form of statement is, we can neither make nor destroy matter we can alter its form only. 54 INORGANIC CHEMISTRY FOR COLLEGES The first expression of the law is too broad, for it transcends our experience; a better way to express it is to state that the total amount of matter in any system undergoing change remains con- stant, and by system we mean the sum total of that particular matter involved in the change. We can now appreciate defi- nitely what is meant by the expression, a law of nature. In science a law is a general statement which expresses the conclusions drawn from the study of related facts. The law of the conserva- tion of matter is a statement drawn from a study of the quantity of matter involved when matter changes its form. It is the result of human experience gained through observation and experi- ment. 61. Law of Definite Proportions. As soon as the importance of the quantitative aspect of chemistry was recognized, scientists undertook a study of the quantitative composition of chemical compounds. After exhaustive investigations and extended con- troversies with other chemists Proust, a French scientist, estab- lished facts which warranted the conclusion that the composition of any definite substance remains constant. This conclusion is expressed in the law of definite proportions, which states that any chemical compound contains the elements of which it is com- posed in a definite and unvarying proportion by weight. As applied to a particular example this law says that the composition of the substance water is fixed; it always contains the element hydrogen and the element oxygen, and the proportion by weight in which these elements are present never varies; all samples of pure water under all conditions consist of one-ninth hydrogen and eight-ninths oxygen. 62. Law of Multiple Proportions. It has been stated that there are two oxides of hydrogen water and hydrogen peroxide and that the existence of these two compounds is the result of the fact that hydrogen and oxygen can unite in different proportions by weight. Similar relations exist in the case of other elements; there are two important oxides of carbon, carbon monoxide and carbon dioxide; there are three well-characterized oxides of iron; and the element nitrogen unites with oxygen in five different proportions. The English chemist John Dalton, whose name is one of the most famous in the history of chemistry, was studying such facts as these when he made a very important discovery. THE ATOMIC THEORY. CHEMICAL EQUATIONS 55 He determined the different proportions by weight in which two elements united and discovered a very simple relation between them. He found that in carbon monoxide the carbon and oxygen were united in the relation of 12 parts by weight of the former to 16 parts by weight of the latter; in carbon dioxide he found 12 parts of carbon united with 32 parts of oxygen. It wa a striking fact that 12 parts of carbon unite with 16 parts of oxygen to form one compound, and with just twice as much oxygen to form the other. Dalton analyzed other compounds and found that the figures obtained led to similar results. He determined the differ- ent amounts of an element which united with a fixed weight of another element to form the various compounds of the two that could be made. The ratio between the varying amounts was not always one to two as in the case of the two oxides of carbon, but the ratio was always that of simple whole numbers. Many cases have been examined since Dalton published his results. Sulphur forms two oxides; in one the sulphur and oxygen are present in the ratio of 32 parts by weight of sulphur to 32 parts by weight of oxygen; in the other the proportion is 32 of sulphur to 48 of oxygen. The weights of oxygen which unite with 32 parts of sulphur are 32 and 48; these numbers are in the ratio of 2 to 3. There are five oxides of nitrogen. If we determine the weights of oxygen which unite with 28 grams of nitrogen in each case to form these oxides, we shall find that these weights are, respectively, 16, 32, 48, 64, 80 numbers which are in the ratio of 1, 2, 3, 4, 5. On what was perhaps insufficient evidence Dalton enunciated what he called the law of multiple proportions, but subsequent investigation has shown that the law summarizes the facts. The law of multiple proportions states that when two elements unite to form more than one compound, the weights of one of these elements which unite with a fixed amount of the second, are in the ratio of small whole numbers. 63. The Atomic Theory. Dalton was much interested in the changes that take place in the weather; he made daily observa- tions and recorded them for many years. He observed that the atmosphere contained varying amounts of moisture from day to day, and he sought to explain how water-vapor could mix with air. Dalton came to the conclusion that this would be possible if air were composed of very small particles which moved about 56 INORGANIC CHEMISTRY FOR COLLEGES freely and were separated by spaces of considerable size as com- pared with the size of the particles. He considered water-vapor as made up in the same way. If air and water-vapor are brought together the particles can mingle and form a uniform mixture. He called these small invisible particles atoms. When the law of multiple proportions was discovered, Dalton saw at once that his conception of atoms offered a simple explanation of the remarkable facts summarized in the law. The law of definite proportions could also be explained. Dalton stated that all substances are composed of atoms, and that when these unite compounds are formed. Elements are made up of atoms, all of which are alike in substance and weight. Dalton recognized the importance of determining the characteristic weights of the atoms of the several elements, but as it was impossible to isolate a single atom and weigh it, he took upon himself the task of determining their rela- tive weights. We shall see later how this can be done. Dalton proposed the atomic theory in a book entitled " A New System of Chemical Philosophy," which was published in 1808. 64. Let us take the present-day conception of atoms as developed from the theory of Dalton and see how it furnishes an explanation of the quantitative relationships observed in chemical action. All substances are made up of exceedingly small particles called atoms. The atoms of any particular element are alike in substance and weight, but differ in these two respects from the atoms of all other elements. Chemical combination consists in the union of two or more atoms to form what is called a molecule. The smallest particle of a chemical compound is a molecule, since chemical compounds contain two or more elements. This comparatively simple conception of the composition of matter has been the guiding principle in chemistry for over a hundred years. While at first it was a mere hypothesis a reasonable guess which was put forward to help explain a few important facts, it has grown in importance as it has been investigated and has been applied to new discoveries made as the science grew. It has become so firmly established as the basis of all science that it is recognized to-day as a fact. Experiments of great brilliancy in conception and ingenuity in execution have demonstrated the presence of atoms, which have been weighed and counted. When we recognize the fact that the amount of hydrogen that can be THE ATOMIC THEORY. CHEMICAL EQUATIONS 57 held in a thimble contains about the number of atoms expressed by 2 followed by 20 ciphers, we can begin to appreciate to what extent man can go in unfolding what used to be called the secrets of nature. The atomic conception of matter furnishes a reasonable expla- nation of the facts summarized in the law of definite proportions and the law of multiple proportions. 65. Fact, Law, Hypothesis, Theory, Science. In the growth of knowledge facts are first discovered through observation and experiment. These facts are next sorted out, and those of the same nature are grouped together and their relation one to the other studied. If a general statement can be made which sum- marizes a large number of facts, a law is discovered. But inquiry does not stop here; thinking men endeavor to find some reason for these laws; they picture to themselves how the fundamental concepts, matter and energy, can produce the observed phenomena; and the result is a hypothesis. The hypothesis is now scruti- nized carefully; is it in accord not only with the facts and laws which led to its production, but with all known facts? The proc- ess known as deductive reasoning is next applied; if the hypothe- sis is true what consequences follow? Can we foretell any undis- covered facts? The conclusions which have been deduced are next tested; if they prove not to be in accord with the facts the hypothesis is rejected or modified. When a hypothesis has been repeatedly tested in this way, and has been shown to be in accord with all known facts it is called a theory. This way of increasing knowledge is called the scientific method. The sum-total of the facts, laws, and theories pertaining to any particular branch of knowledge is called a science. Science has been defined as sys- tematized human knowledge. The development of any particular science is dependent upon the extent to which the facts considered in that science have been systematized. Some sciences are largely descriptive, that is, only a few great generalizations have been drawn from the facts; others are highly developed. The more the processes of mathematics can be used to summarize the facts of a science, the more developed becomes that science. Physics is highly developed; we shall see later, for example, that the behavior of gases under varying conditions of volume, pressure, and temperature can be expressed by a very simple mathematical 58 INORGANIC CHEMISTRY FOR COLLEGES equation; thousands of isolated facts are indicated by the expres- sion pv = RT, when the significance of the letters used in the equation is understood. Chemistry was for a long time a descrip- tive science, but its development in recent years has been rapid. 66. Atomic Weights. We have learned that when two ele- ments unite to form a compound the relation between the weights of the two is definite. These relations have been determined with great care. It has been found, for example, that 16 parts by weight of oxygen unite with 2 parts by weight of hydrogen, 16 of sulphur, 12 of carbon, 65 of zinc, and so forth. Applying the theory of atoms and molecules to such figures as these, it has been possible to assign numbers to the various atoms which repre- sent their relative weights; the number assigned to each element is called its atomic weight. The interesting way in which this is done will be described later, since it is necessary to acquire a more extensive knowledge of chemistry before it can be thoroughly understood. Hydrogen is the lightest substance known and its atom was taken, therefore, as the standard of weight; the atomic weight of hydrogen is 1. Since an atom of oxygen weighs sixteen times as much as an atom of hydrogen, its atomic weight is 16. Similarly when we say the atomic weight of carbon is 12 we mean that an atom of carbon is 12 times as heavy as an atom of hydrogen. A list of the elements with their atomic weights is given on the inside of the back cover of this book. The student will have occasion to use these numbers repeatedly, but it is inadvisable to attempt to learn them by heart; as the more important ones are used over and over again, their values will be remembered without effort. 67. Symbols and Formulas. The use of letters to represent atoms has materially simplified the method of recording chemical reactions. The alchemists used astronomical figures to represent the elements; the symbol O or >|< for the sun was used to repre- sent gold, the crescent moon C f r silver, etc. Dalton employed geometrical figures; O represented hydrogen, O oxygen, car- bon, O O carbon dioxide, etc. It was some years later, in 1811, when the practical suggestion was made by Berzelius to indicate the elements by letters taken from their names. stands for oxygen, H for hydrogen, N for nitrogen, etc. When THE ATOMIC THEORY. CHEMICAL EQUATIONS 59 the initial letter of two or more elements is the same, a second letter taken from the name is added in some cases; thus, C is the symbol for carbon, Ca for calcium, and Cd for cadmium. In the case of some of the metals the symbol is taken from the Latin name of the element; Sn is the symbol for tin and is derived from the word stannum; ferrum, iron, gives us the symbol Fe for this metal, etc. 68. The composition of a compound is represented by a chem- ical formula, which is made up of the symbols of the elements present in the compound. Used for this purpose the symbol signifies one atom of the element, or that weight of the element which is numerically equal to its atomic weight. We can use any system of weights, but if a statement is not made to the contrary the gram is the unit understood. The atomic weight of carbon, for which the symbol C is used, is 12; the atomic weight of oxygen is 16 and its symbol is 0. When these symbols appear in a chemi- cal formula, as CO, it means that the molecule of the substances represented by the formula contains one carbon atom and one oxygen atom; since these atoms have definite weights, twelve and sixteen respectively, the formula indicates that in this sub- stance the elements are united in the proportion of 12 parts by weight of carbon to 16 parts by weight of oxygen. The molecular weight of a compound is the sum of the weights of the atoms it con- tains. If grams are used in any calculation involving the above formula, the number represents what is called a gram-molecular- weight or mol, that is, 12 + 16 = 28 grams of the compound. Likewise 28 pounds is a pound-molecular-weight of it. 69. The substance to which is assigned the formula CO is called carbon monoxide to distinguish it from another oxide of carbon the molecule of which is made up of one carbon atom and two oxygen atoms carbon dioxide. The prefixes indicate in these and other cases the number of atoms of any particular ele- ment present; they are derived from Latin words signifying one, two, three, etc. When more than one atom of an element is present in a com- pound, this fact is indicated in its formula by writing a number to the right and below the symbol of that element; the formula for carbon dioxide, which contains one carbon and two oxygen atoms, is expressed thus: C02. A number placed in front of a formula 60 INORGANIC CHEMISTRY FOR COLLEGES indicates that a certain number of these molecules is considered; thus 2CO2 stands for 2 molecules of carbon dioxide. 70. Chemical reactions can be clearly represented by means of chemical formulas. For reasons which will be given later we believe that oxygen gas is made up of molecules each of which contains 2 atoms. Atoms are represented by symbols; mole- cules, which contain 2 or more alike or unlike atoms, are repre- sented by formulas; the formula for oxygen gas is, accordingly, 62. Likewise the formula for hydrogen gas is H2. Water has been shown to be made up of 2 hydrogen atoms and 1 oxygen atom; its formula is, accordingly, H^O. With the aid of these formulas we can express the fact that oxygen unites with hydro- gen to form water; it is represented as follows, by what is called a chemical equation: 2H 2 + O 2 = 2H 2 This equation signifies that 2 molecules of hydrogen unite with 1 molecule of oxygen and form 2 molecules of water. Defi- nite weight relations are also indicated. Since hydrogen is the lightest substance known the weight of its atom may be taken as the unit of weight of atoms. A hydrogen atom weighs 1 unit and an oxygen atom weighs 16; consequently, a molecule of hydrogen, H2, weighs 2, and one of oxygen, 62, weighs 32. A molecule of water weighs 2 + 16 = 18. The equation states that 2 mole- cules of hydrogen weighing 2X2=4, unite with 1 molecule of oxygen weighing 32 and form 2 molecules of water weighing 2 X 18 = 36. To express this clearly we can rewrite the equation as follows: 2H 2 + O 2 2H 2 O 2(1 + 1) (16 + 16) 2(1 + 1 + 16) 4 + 32 =36 These numbers are relative weights; we can use them in any units. f grams j f grams 1 Four j pounds ! of hydrogen react with 32 \ pounds > of oxygen and I ounces J I ounces J (grams ^ pounds \ of water, ounces J THE ATOMIC THEORY. CHEMICAL EQUATIONS 61 When symbols are put together in the way shown above the result is called an equation, because the number of atoms on one side of the equality sign equals the number on the other; the equation has a quantitative significance. We could express the fact that hydrogen and oxygen united to form water, in the fol- lowing way: _ H 2 + O 2 The arrow signifies that the substances represented by the formulas to the left of it change to the substance indicated by the formula at its right. This qualitative expression, signifying only the sub- stances involved, lacks the definiteness and fullness of an equation, which furnishes both qualitative and quantitative information. Some authors do not use the equality sign in writing chemical equations; they use an arrow for all purposes to indicate a chemi- cal change. In this book the equality sign and arrow will have the significance which has just been explained. 71. The Writing of Chemical Equations. In order to write a chemical equation we first set down the formulas of the substances which interact, separating these by plus signs to show clearly each formula. We next write in the same way the formulas of the products formed, and indicate by the use of the arrow that a transformation has taken place. The reaction between zinc and hydrochloric acid is expressed in this way as follows : Zn + HC1 -> ZnCl 2 + H 2 It is necessary to learn the formulas of substances as we meet them. Formulas express facts which have been established as the result of experiment. Zinc chloride contains the metal and chlorine in the proportions represented by the formula; and it is necessary to remember that the formula of zinc chloride is ZnCl 2 . We shall soon see that if we know this formula and a few other facts we can readily write the formulas of a large number of com- pounds which contain zinc. The task of remembering many formulas is not so great as might appear at first; there is a beau- tiful system underlying chemical combination, which once learned, removes the necessity of relying too often on an act of mere memory. Returning to the chemical equation under discussion, we see 62 INORGANIC CHEMISTRY FOR COLLEGES that what has been written does not express definitely the quanti- ties involved in the reaction. We must next modify the expres- sion to include this; we balance the equation, as chemists say. If an equality sign is to be put between the formulas of the react- ing substances and those formed, the resulting equation must not violate the law of the conservation of matter; for example, we must represent as much hydrogen on one side of the sign as on the other. We next examine each symbol and see if the same number of atoms appear on the two sides of the arrow. There is 1 zinc atom on the left, and 1 on the right; no change is required here. There is 1 hydrogen atom on the left and there are 2 on the right; a modification is necessary. If we take 2 molecules of hydrochloric acid, 2HC1, we have 2 hydrogen atoms; this is accordingly done and a 2 is placed in front of the formula; the expression now becomes, Zn + 2HC1 -> ZnCl 2 + H 2 We continue the inspection and find that there are 2 chlorine atoms on either side of the arrow. The number of atoms of each element on one side of the arrow appears to be the same as that of the same element on the other side. We run through the inspec- tion of each element a second time to see that the placing of the 2 in front of HC1 has not altered the relations which were examined before this change was made, and if we find that we have the same number of zinc atoms, the same number of hydrogen atoms, and the same number of chlorine atoms on the two sides of the arrow, we conclude that the formulas have been properly balanced; we change the arrow to an equality sign, and have, then, a chemical equation : Zn + 2HC1 = ZnCl 2 + H 2 A very simple example has been taken to show how equations are balanced; the consideration of a more difficult one will be in- structive. When steam is passed over hot iron, hydrogen and an oxide of iron are formed. Experiment has shown that the oxide has the formula FesCU; this is a fact which has to be remembered. If we know the formula of water and that of hydrogen we can write the equation for the reaction. We shall first set down the symbol of iron and the formula of water, since these two substances THE ATOMIC THEORY. CHEMICAL EQUATIONS 63 interact, then write an arrow, and next the formulas of the sub- stances formed, thus: Fe + H 2 O -> Fe 3 O 4 + H 2 Next we must balance; we see 1 iron atom on the left, and 3 on the right; place a 3 before the symbol of iron; the expression is now as follows: 3Fe + H 2 O -* Fe 3 O 4 + H 2 If the iron is to stay balanced we cannot change the number of molecules of Fe 3 4 ; we therefore next balance the oxygen, the other element in the compound containing the atoms just balanced. There are 4 atoms of oxygen on the right of the arrow; there must be 4 on the left, so we take 4 molecules of water. This changes our expression to the following: 3Fe + 4H 2 - Fe 3 O 4 + H 2 The number of molecules of water must not be changed now, for if we do this, the balancing of the iron will be destroyed; we therefore balance hydrogen next. There are 8 atoms of this element represented to the left of the arrow; there must be 8 on the other side; and so a figure 4 is put in front of the formula for hydrogen. The expression now becomes, 3Fe + 4H 2 O -> Fe 3 O 4 + 4H 2 It is next studied to see if all the atoms are balanced; this proves to be correct, and when the arrow is replaced by the sign of equal- ity we have a correct chemical equation. The advisability of checking up the work by a re-examination of the final equation cannot be impressed too strongly, for such an examination will detect errors that might have been made. Another point is to be strongly emphasized: In balancing equations the numbers used in the formulas of substances to express their composition, such as the 3 and the 4 in Fe 3 4 , can never be changed. As has been pointed out, these numbers indicate the relative proportions of .the elements present in the compounds. For example, the oxide of iron formed by the action of steam on iron always has the composition indicated by the formula Fe 3 4 it contains iron and oxygen in the proportion of 3 X 56 parts by weight of iron to 64 INORGANIC CHEMISTRY FOR COLLEGES 4 X 16 parts by weight of oxygen, 56 and 16 being the atomic weights of iron and oxygen respectively. Again, if in balancing an equation containing the formula H2O, which represents the substance water, we change the formula to H 2 (>4 to bring the oxygen into balance, we would make a great mistake, for there is no compound with the formula H 2 (>4, and if there were it would not be water. To state the fact in other words, in balancing equations all we can do is to put what numbers we please before the formulas and symbols. This change means that we use and obtain as much of each substance as is required by the law of the conservation of matter. This subject has been gone into at such length on account of its importance, and the fact that the student is apt to make the mistake which he is here told to avoid. When a symbol appears in the formulas of two substances on the same side of an equation, balancing is, at times, a little more difficult. We will write the equation for the reaction which takes place between sodium peroxide and water, which it will be recalled, yields sodium hydroxide and oxygen. In order to be able to do this it is necessary to know the formulas of all the sub- stances involved; these are given below: 1 Na 2 O 2 + H 2 O - NaOH + 2 We start by balancing one element, usually the first written down, which is sodium in this case. Proceeding in the way illustrated above we get the expression, Na 2 O 2 + H 2 O -> 2NaOH + O 2 Having placed a 2 in front of the NaOH we next balance the oxygen, an element present in this compound. As it now stands there are 3 oxygen atoms to the left of the arrow and 4 to the right; the oxygen atoms can be balanced by taking 2 mole- cules of water: Na 2 O 2 + 2H 2 O -> 2NaOH + O 2 We have just changed the number of water molecules to balance oxygen, so we now attempt to balance the hydrogen, the second element in water. There are 4 atoms to the left of the arrow and 2 to the right. To balance now we must change the number 1 The symbol for sodium (natrium) is Na. THE ATOMIC THEORY. CHEMICAL EQUATIONS 65 in front of the formula for sodium hydroxide from 2 to 4, and this throws out of balance the sodium and perhaps the oxygen. It is evident that the equations cannot be balanced if we use but one molecule of Na2O2. We accordingly place a 4 in front of the formula NaOH, and balance the sodium again by putting 2 before Na 2 O 2 . Starting with 2Na2O 2 we go through the entire process again and find this time that when we reach the last element it balances. The various steps, which should be analyzed carefully by the student, are as follows: Na 2 2 + H 2 O - NaOH + O 2 2Na 2 O 2 + H 2 O - 4NaOH + O 2 2Na 2 O 2 + 2H 2 O = 4NaOH + O 2 The student should solve the problems in balancing chemical equations given at the end of the chapter. The method is used constantly and relieves the memory of much effort; it is not neces- sary to remember the number of molecules entering into a reaction. The formulas of the reacting substances and those formed should be kept in mind; this soon becomes an easy task, since these formulas occur again and again in other connections. 72. Some Simple Chemical Equations. It will be instructive to express in equations the chemical changes which have been studied thus far. Magnesium burns in oxygen to form magnesium oxide: 2Mg + O 2 = 2MgO When mercury 1 and iodine are rubbed together mercuric iodide is produced: Hg + I 2 = HgI 2 Iron and sulphur when heated are converted into ferrous sul- phide: Fe + S = FeS 1 The symbol for mercury is Hg; it is derived from the Latin name of the metal, hydrargyrum, which means liquid silver. The symbol for iron, Fe, comes also from the Latin name of the metal which is ferrum; K is the symbol for potassium, which was called kalium. 66 INORGANIC CHEMISTRY FOR COLLEGES Oxygen can be prepared by heating mercuric oxide: 2HgO = 2Hg + 2 by heating potassium chlorate : 2KC1O 3 = 2KC1 + 3O 2 by the electrolysis of water : 2H 2 O = 2H 2 + O 2 and by the action of water on sodium peroxide: 2Na 2 O 2 + 2H 2 O = 4NaOH + O 2 The chemical conduct of oxygen was shown by burning in it carbon, sulphur, phosphorus, and iron. The equations for these reactions are as follows: C + 2 = C0 2 S + O 2 = S0 2 4P + 5O 2 = 2P 2 O 5 3Fe + 2O 2 = Fe 3 O 4 Hydrogen was prepared by the electrolysis of water: 2H 2 O = 2H 2 + O 2 It was formed by the reaction between water and sodium: 2Na + 2H 2 - 2NaOH + H 2 Iron reacted with steam to produce hydrogen: 3Fe + 4H 2 O = Fe 3 O 4 + 4H 2 Hydrogen was also prepared by the action of certain metals with hydrochloric acid, HC1, and sulphuric acid, H 2 SO4,* the equa- tions for these reactions are as follows: Zn + 2HC1 = ZnCl 2 + H 2 Fe + 2HC1 = FeCl 2 + H 2 Sn l + 2HC1 = SnCl 2 + H 2 Zn + H 2 SO 4 = ZnSO 4 + H 2 Fe + H 2 SO 4 = FeSO 4 + H 2 1 Sn is the symbol for tin; it is derived from the Latin word stannum. THE ATOMIC THEORY. CHEMICAL EQUATIONS 67 Hydrogen is formed when zinc and aluminium react with sodium hydroxide: Zn + 2NaOH = (NaO) 2 Zn + H 2 2A1 + 6NaOH = 2(NaO) 3 Al + 3H 2 Hydrogen burns in oxygen: 2H 2 + O 2 = 2H 2 O and reduces hot copper oxide: CuO + H 2 = Cu + H 2 O EXERCISES 1. Carbon monoxide contains 57.14 per cent oxygen, (a) How many grams of oxygen are there in 100 grams of carbon 'monoxide? (6) How many grams of carbon in 100 grams of carbon monoxide? (c) How many grams of oxygen are combined with 1 gram of carbon in carbon monoxide? Carbon dioxide contains 72.77 per cent of oxygen, (d) How many grams of oxygen are combined with 1 gram of carbon in carbon dioxide? (e] What is the relation between the weight of oxygen combined with 1 gram of carbon in carbon monoxide and carbon dioxide. 2. There are two oxides of sulphur, one of which contains 50 per cent of oxygen and the other 60 per cent, (a) Calculate the weight of oxygen united with 100 grams of sulphur in each case. (6) What relation do these numbers bear to each other? (c) Is the result in accord with the law of multiple proportions? 3. Convert the following into balanced chemical equations: (a) Al + HC1 -* A1C1 3 + H 2 . (6) Mn + HC1 -* MnCl 2 + H 2 . (c) Sb + O 2 -* Sb 2 O 3 . (d) Fe 2 O 3 + H 2 - Fe + H 2 O. (e) Ca + H 2 O -> Ca(OH) 2 + H 2 . (/) Na 2 O 2 + H 2 SO 4 -t Na 2 SO 4 + H 2 O 2 . (g} MnO 2 - Mn 3 O 4 + O 2 . (h) CuO - Cu 2 O + O 2 . CHAPTER VII CHEMICAL CALCULATIONS 73. It was shown in the last chapter that the conception of atoms and molecules and a study of the quantitative relations in chemical change led to a simple method of expressing in the form of equations the reactions which take place between substances. Since these equations express so many facts in a very brief form, they are constantly used in chemistry, and it is essential that their significance and the uses to which they can be put are thoroughly understood. Let us take a simple chemical reaction and see how an equation was written for it and how this equation can be used when once it is known. When hydrogen is passed over hot copper oxide the products formed are copper and water. If the atomic weights of the elements are known, a study of the quantitative relations involved will give the information necessary to write the equation. The atomic weights of copper, hydrogen, and oxygen are 63, 1, and 16 respectively. Experiments showed that 63 grams of copper unite with 16 grams of oxygen to form 79 grams of copper oxide, which, consequently, has the formula CuO. Experiments also showed that 2 grams of hydrogen unite with 16 grams of oxygen to form 18 grams of water; the formula of the latter must be, therefore, IbO. The quantitative relations found in the reaction between copper oxide and hydrogen are as follows : 79 grams of copper oxide react with 2 grams of hydrogen to form 63 grams of copper and 18 grams of water. These facts are evidently repre- sented by the following equation: CuO + H 2 = Cu -f H 2 O It is important for the student to see clearly how the weights given lead to the conclusion expressed in the equation. The 68 CHEMICAL CALCULATIONS 69 results of the several experiments have been recorded in this equation, and we can now use it to calculate the weight relations involved in the reaction. It is important to note that all chemi- cal equations are based on the previous study of the weights of the substances entering into the reaction; we cannot make them up out of our heads, or guess what happens. Chemical formulas and equations are a kind of short-hand with which we can record in brief form many facts. All the facts in regard to the reaction between copper oxide and hydrogen stated above are recorded in the equation given ; other facts, not yet brought out, are indicated also. When one 'learns how to read this chemical short-hand, every equation is a source of much information. 74. We can now understand why we are able to use chemical equations to calculate the weight relations involved in the change expressed by the equation. Let us see how this is done. How many grams of hydrogen will be required to react with 100 grams of copper oxide? We will rewrite the equation for the reaction and place under the formulas the weight relations shown by the atoms present. We must, of course, know the atomic weights, and we obtain them by referring to the table on the inside of the back cover of this book. Accurate values of the atomic weights are used in these calculations. In order to make the discussion easier to follow, round numbers have been used, heretofore. The accuracy of the result desired determines how many decimal places should be retained in the atomic weights and the calcula- tions. The equation for the reaction described above is as fol- lows: CuO + H 2 = Cu + H 2 (63.57 + 16) + (1.008 + 1.008) 63.57 + (1.008 + 1.008 + 16) 79.57 2.016 63.57 18.016 By proceeding in the way indicated, we discover the relative weights of the several substances which take part in the reaction. From these numbers we can calculate by simple proportion what weight of any substance is involved when the weight of any other is specified. We can answer, for example, the question stated above how many grams of hydrogen will be required to react with 100 grams of copper oxide? The equation tells us that 79.57 grams of the oxide react with 2 grams of hydrogen, then 70 INORGANIC CHEMISTRY FOR COLLEGES CuO : H 2 CuO : H 2 , 79.57 : 2.016 = 100 : x, 79.57z = 201.6 *?ni fi x = ' = 2.53 grams of hydrogen. / y.o i It is advisable to write over the terms of the proportion the formulas of the substances involved; neglect to do this often results in an incorrect statement of the proportion. Another form of stating the proportion is to be recommended; it reduces somewhat the amount of work required. The .chemical equation is written, as before, and above the formulas of the substances involved in the particular problem are placed the given weight and the letter x to indicate the weight sought. Beneath these formulas are written the molecular weights, which are obtained by adding the atomic weights of the atoms in the separate mole- cules. In the case of the problem just solved this procedure would lead to the following expression: 100 gms. x gins. CuO + H 2 = Cu + H 2 O 79.57 2.016 From these numbers we make a simple proportion in the order as written, thus, 100 : x = 79.57 : 2.016 79.57s = 201.6 x = 2.53 grams of hydrogen It is evidently unnecessary to make the calculations of the molec- ular weights of the substances not involved in the particular problem being solved. In order to illustrate the method further, the solutions of other problems based upon this reaction will be indicated. How many grams of water are produced when 25 grams of copper oxide are reduced by hydrogen? 25 gms. x gms. CuO + H 2 = Cu + H 2 79.57 18.016 25 :x = 79.57:18.016 79.57x = 25 X 18.016 x = 5.66 grams of water CHEMICAL CALCULATIONS 71 How many grams of copper can be obtained by reducing 5 grams of copper oxide by hydrogen? 5 gms. x gms. CuO + H 2 - Cu + H 2 O 79.57 63.57 5 : x = 79.57 : 63.57 79.57z = 63.57 X 5 x = 3.99 grams of copper In this case we can simplify the solution somewhat. As all the copper comes from the copper oxide we can set down the problem as follows: 5 : x CuO : Cu 79.57 63.57 The rest of the solution is as given above. As a further example of the method one more problem will be solved. How many grams of oxygen can be obtained by decom- posing 10 grams of sodium peroxide with water? 10 gms. x gms. 2Na 2 2 + 2H 2 O = 4NaOH + O 2 2(23 + 23 + 16 + 16) 16 + 16 156 32 10 : x = 156 : 32 156z = 10 X 32 x = 2.05 grams of oxygen. In the above equation 2 molecules of sodium peroxide take part in the reaction. Each molecule contains 2 sodium atoms each weighing 23, and 2 oxygen atoms each weighing 16; the molecule weighs, accordingly, 23 + 23 + 16 + 16 = 78. Since 2 mole- cules are involved, this number must be multiplied by 2. 75. The same method of calculation can be used to determine the percentage of an element in a compound the formula of which is known. The words per cent are a contraction of per centum; 10 per cent means ten per hundred. If we inquire what is the percentage of oxygen in copper oxide, the question is answered by 72 INORGANIC CHEMISTRY FOR COLLEGES stating how many grams of oxygen are contained in 100 grams of copper oxide. This can be determined as follows: 100 gms. CuO 63.57 + 16 79.57 100: x 79.57z X x gms. : 16 = 79.57 : 16 = 1600 = 20.09 per cent. What percentage of potassium chlorate is oxygen? 100 gms. x gms. KC1O 3 : 3O 39.10 + 35.46 +16 + 16+16 16 + 16 + 16 122.56 48 100 : x = 122.56 : 48 122.56z = 48 X 100 x = 39.16 per cent 76. Calculations Involving Volumes of Gases. It is often necessary to know the volume of a gas formed in a given reaction. This can be done with the aid of the chemical equation for the reaction and the knowledge of the weight of a liter of the gas. The weight of the gas formed is first determined in the way just described, and then the volume of this weight calculated. A study of the formulas of gases from the standpoint of volumes leads to a result that makes the method of calculation simpler and does not involve a knowledge of the weight of any particular gas. Since in all equations the molecular formulas of the gases are used and these represent different weights, for example, Eb represents 2 grams of hydrogen and 02 32 grams of oxygen, it would be well to know what volumes of the several gases these formulas represent. These can be readily calculated. What is the volume of 2 grams of hydrogen? One liter of the gas has been found by experiment to weigh 0.09 gram at and 760 mm. If 1 liter weighs 0.09 gram how many liters weigh 2 grams? 1 : 0.09 = x : 2, x = 22 A liters CHEMICAL CALCULATIONS 73 The formula for oxygen is 62 and its molecular weight is 2 x 16 = 32. What is the volume of 32 grams of oxygen? One liter of oxygen weighs 1.429 grams at and 760 mm.: 1 : 1.429 = x : 32, x = 22 .4 liters If we calculate in this way the volume of a gram-molecular-weight (68) of any gas the result is always 22.4 liters. We shall learn later (Chapter XXIV) the reason for this striking fact. We see, then, that the formulas of gases have a definite significance in regard to volume as well as to weight. The formula CO2 assigned to carbon dioxide means, as we have learned, that the molecule of the gas contains 1 atom of carbon and 2 atoms of oxygen, and since the atomic weight of carbon is 12 and that of oxygen 16, it means, further, that 12 grams of carbon are united with 2 X 16 =32 grams of oxygen. We see now that 44 grams of CO2 (12 + 2 X 16= 44), which is 1 gram-molecular-weight, has a volume of 22.4 liters. We are now in a position to calculate the volumes of gases found in any chemical reaction. The method is illustrated by the following example. What volume of hydrogen is formed when 10 grams of zinc dissolve in hydrochloric acid? We write the equation as before, 10 gms. x liters Zn + 2HC1 = ZnCl 2 + H 2 65.37 22.4 liters, but place below the formula for hydrogen the volume of the gas formed and use this in the proportion instead of the weight in solving the problem, 10 : x = 65.37 : 22.4 liters x = 3.42 liters. In the following problem the question is put in a different way but the method is the same. What weight of zinc is required to liberate 5 liters of hydrogen? x gms. 5 liters Zn + 2HC1 = ZnCl 2 + H 2 65.37 22.4 liters x : 5 = 65.37 : 22.4 x = 14.6 grams. 74 INORGANIC CHEMISTRY FOR COLLEGES In solving problems in this way the volume is always 22.4 liters for each gram-molecular-weight whatever the gas may be. If 3 molecules, for example, appear in the equation, then the volume of the gas is 3 times 22.4 liters. In the case of the decomposition of potassium chlorate the equation is as follows: 2KC1O 3 = 2KC1 + 3O 2 2[39 + 35 + 3(16)] 3(22.4 liters) 244 67.2 liters 77. The volume of 1 gram-molecular-weight of a gas, namely, 22.4 liters, is called the gram-molecular-volume. Since this is a constant quantity, it is possible to see at a glance the relation by volume in which gases interact, if the equation for the reaction is known. In the case of the union of hydrogen with oxygen to form water: 2H 2 + O 2 = 2H 2 O 2(22.4 liters) 22.4 liters 2(22.4 liters) if the temperature is such that the water formed is a gas i.e., steam we see that 2 times 22.4 liters of hydrogen unite with 22.4 liters of oxygen to form 2 times 22.4 liters of steam, which means that 2 volumes of hydrogen react with 1 volume of oxygen to form 2 volumes of steam. The relation between the volumes of the gases involved in the reaction is the same as that between the number of molecules. If we take as a second example the equation, N 2 + 3H 2 = 2NH 3 which expresses the union of nitrogen with hydrogen to form ammonia, which is a gas, we see that 1 molecule of nitrogen unites with 3 molecules of hydrogen to form 2 molecules of ammonia and that 1 volume of nitrogen unites with 3 volumes of hydrogen to form 2 volumes of ammonia. The relation holds, of course, whatever unit of volume is used liters, pints, gallons, bar- rels, etc. Since 1 gram-molecular-weight of all gases occupies 22.4 liters we can calculate from the formula of a gas what 1 liter of it weighs. For example, the formula of sulphur dioxide is S0 2 . Since the atomic weight of sulphur is 32 and that of oxygen 16, the molecu- ular weight of sulphur dioxide is 32 + (2 X 16) = 64, Sixty- CHEMICAL CALCULATIONS 75 four grams of the gas have the volume 22.4 liters, therefore, 1 liter weighs 64 grams -5- 22.4 = 2.85 grams. 78. By making use of the fact that the volumes of 1 gram- molecular-weight of all gases are the same, we can readily tell by inspection of the formulas of two gases which is the heavier. Which is heavier, nitrogen or oxygen? The molecular weight of the former, N 2 , is 2 X 14 = 28, and of the latter, O 2 , 2 X 16 = 32. The weight of 22.4 liters of nitrogen is 28 grams and of oxygen 32 grams; the latter is, accordingly, the heavier gas. Which is the heavier gas, carbon dioxide, CO2, or sulphur dioxide, 862? The molecular weight of the former is 12 + (2 X 16) = 44, and of the latter 32 + (2 X 16) = 64; sulphur dioxide is heavier. Pure dry air is essentially a mixture of oxygen and nitrogen molecules. The two gases are present in such proportions that the average weight of the molecules is 28.8; any gas the molecular weight of which is greater than 28.8 is heavier than air. 79. In commercial work weights are usually expressed in pounds and the volume of gases in cubic feet. It can be readily calculated that if 1 gram-molecular-weight of a gas has a volume of 22.4 liters, 1 pound-molecular-weight has a volume of 359 cubic feet. The latter relation is commonly made use of in the calculations involved in technical work. EXERCISES NOTE: The atomic weight of the elements can be found in the table on the inside of the back cover. All questions in regard to volume relations refer to standard conditions, and 760 mm. 1. Calculate the molecular weight and the percentage of oxygen in each of the following: (a) H 2 SO 4 , (6) KMnO 4 , (c) P 2 O 6 . 2. Calculate the percentage of each element in potassium chlorate, KC1O 3 . 3. How many grams of hydrogen can be obtained when 25 grams of zinc are dissolved in hydrochloric acid? Zn + 2HC1 = ZnCl 2 + H 2 . 4. How much zinc must be dissolved in sulphuric acid to produce 40 liters of hydrogen? 5. What weight of potassium chlorate is required to furnish enough oxygen to fill five 250 c.c. bottles, assuming that one-quarter of the gas evolved is lost in collecting the gas. 6. (a) What weight of carbon is required to react with 100 pounds of zinc oxide according to the folio \ving equation: ZnO + C = Zn + CO? (6) How much zinc is formed? (c) What is the weight of the carbon mon- oxide formed and what is its volume in cubic feet? 76 INORGANIC CHEMISTRY FOR COLLEGES 7. It was found by experiment that 1 liter of chlorine, C1 2 , weighs 3.17 grams. Calculate the volume of 1 gram-molecular-weight of the gas. 8. Calculate the weight of 1 liter and 1 cubic foot of each of the following gases: (a) HC1, (6) SO 2 , (c) N 2 , (d) CO, (e) CO 2 . 9. Arrange the gases having the following formulas in the order of increas- ing density: N 2 , CO, NO, CO 2 , SO 2 , NH 3 , CH 4 . 10. Calculate the weight of 1 liter of a mixture containing (a) equal volumes of oxygen and nitrogen, (b) 1 volume of oxygen and 4 volumes of nitrogen. 11. (a) Calculate the number of cubic feet of air required to burn 1 long ton of coke which contains 87 per cent of carbon : C -f- O 2 = CO 2 . (6) What is the relation between the volume of the air used and the volume of gases issuing from the furnace after they have been cooled to the original temperature of the air? (c) If the carbon burns to carbon monoxide 2C + O 2 = 2CO what is the relation between the volumes asked for in (b) above? 12. Calculate the volume in cubic feet of 1 pound-molecular-weight if 1 gram-molecular-weight has a volume of 22.4 liters. 13. An experiment showed that when 9.75 grams of zinc were dissolved in acid, 0.15 gram of hydrogen was formed. Calculate the atomic weight of zinc taking 1 as the atomic weight of hydrogen. 14. It was found in an experiment that 2 grams of silver united with 0.657 gram of chlorine to form silver chloride: 2Ag + C1 2 = 2AgCl. Assum- ing the atomic weight of chlorine as 35.46, calculate the atomic weight of silver. CHAPTER VIII MEASUREMENT OF GASES 80. On account of the fact that gases are so relatively light it is very difficult to weigh them with any degree of accuracy. The amount of hydrogen which weighs as much as a five-cent piece has a volume under ordinary conditions of over 57 liters or 15 gallons. The accuracy of the results obtained in weighing such bulky substances is markedly affected by the fact that the weight of the gas is very small compared with the weight of the vessel which contains it. The greatest experimental skill and the most refined apparatus are necessary to obtain accurate results in weighing gases. We can, however, measure easily the volume a gas occupies, and since the relations between the weights and volumes of all gases have been determined with great care, the weight of any sample of gas the volume of which is known, can be determined by calculation. When gases are heated they expand, and when they are sub- jected to pressure their volume decreases. For this reason if the weight of a definite volume of gas is to be recorded, a statement must be made as to the temperature and the pressure under which the gas was measured. It is not necessary, however, to weigh gases under all possible conditions of temperature and pressure, because simple laws have been discovered in regard to the behavior of gases when these two factors change. 81. Thermometers. Before discussing these important laws it is necessary to get a definite knowledge of how temperature and pressure are measured. Temperature is measured by a ther- mometer, which is an instrument based on the fact that certain substances increase in volume when they are heated. A ther- mometer is made by blowing a small bulb on the end of a long glass tube, the internal diameter of which is very small. The bulb and the lower part of the tube are filled with mercury and 77 78 INORGANIC CHEMISTRY FOR COLLEGES then placed in melting ice. The place on the tube where the sur- face of the mercury stands is marked. The whole apparatus is next put in the steam rising from boiling water; the mercury increases in volume as it is heated and rises in the tube. When the surface of the mercury no longer rises the point where it stands is marked. These two temperatures the melting-point of ice and the boiling-point of water are the fixed points on thermometers. If the thermometer is to be graduated with a Fahrenheit scale, the one commonly used in daily life, the temp- erature of melting ice is marked 32 and that of boiling water 212. The space between these is divided into 180 divisions (212 32) which are called degrees. The stem of the thermometer is marked above the 212 point and below the point with divisions equal in length to those between these two temperatures. It is, there- fore, possible to read temperatures with a mercury thermometer from the freezing-point to the boiling-point of the metal. On the centigrade thermometer the freezing- and boiling-points of water are marked and 100 respectively, and there are 100 degrees between them. On account of its convenience the centi- grade scale is used in all scientific work except that with which the public come directly in contact, such as meteorology (weather observations), steam engineering, etc. Unless it is noted to the contrary all temperatures recorded in this book refer to the centi- grade scale. 82. Measurement of Pressure. The numerical value of any pressure is usually expressed by stating the height of a column of mercury which exerts the same pressure. When we say, for example, that the pressure of the atmosphere is 76 cm., we mean that the air presses down on any given area with the same force that a column of mercury 76 cm. high which covered this area, would exert. A column of mercury 1 cm. high and 1 sq. cm. in area weighs 13.5955 grams. The absolute unit of pressure is that exerted by 1 gram on 1 sq. cm. To express the pressure of the atmosphere in this unit we must accordingly multiply 76 by 13.5955; it is thus 1033.3 grams per square centimeter. Since, however, pressures are conveniently measured by balancing a col- umn of mercury against them, they are generally expressed in centi- meters of mercury. Pressures are expressed in grams per square centimeter in the so-called C. G. S. system, where the centimeter, MEASUREMENT OF GASES 79 gram, and second are the fundamental units. In expressing pres- sure in commercial work, such as the pressure of steam in a boiler, the unit is the pressure exerted by one pound on one square inch. When a larger unit is required, pressure is expressed in atmos- pheres; an atmosphere is 1033.3 grams per square centimeter or 14.7 pounds per square inch. 83. Atmospheric Pressure. All substances have weight. The gases that make up the air are attracted by the earth and press down upon it. The pressure of the atmosphere is equal to 14.7 pounds on each square inch of the earth's surface; this means that if we could weigh a column of air one square inch in area, which extended from the earth to the extreme limit of the atmosphere it would weigh 14.7 pounds. Of course this cannot be done on an ordinary balance, but it can be done if we balance the weight of the atmosphere against a column of mercury. The instrument used to do this is called a barometer. In its simplest form it is represented by Fig. 9. It is made by carefully filling with mercury a long tube closed at one end. The tube is then covered so that no mercury can escape, inverted, and the lower end placed under the surface of mercury contained in a flat dish; the cover is next removed. If the tube is long enough the mercury will fall until the level of it in the tube has reached a definite height; it will then stand at rest. Since no air was allowed to enter the tube there can be nothing in it above the surface of the mercury except the vapor which is given off by the latter. This is so very small in amount that it can be neglected under most circumstances; the tube contains in the upper part what is called a vacuum. What holds the mercury up against the force of gravity which tends to draw it toward the earth? The air presses down on the mercury in the dish; if that in the tube fell the level of the mercury in the dish would have to rise, and the pressure of FIG. 9. 80 INORGANIC CHEMISTRY FOR COLLEGES the air prevents this. Evidently the mercury in the tube just balances the pressure of the atmosphere, and if we measure the former we determine as a result the latter. If such an instrument as that described is examined, day by day, it will be found that the height of the column of mercury varies. A relationship has been discovered between the height of the mercury in a barometer and the weather; when the barometer is high, that is, when the pressure of the air is great, the weather is apt to be fair; when it falls below a certain point, a storm may be expected. Air travels from a place where there is high pressure, that is, where the air is heavy, to a place where the pressure is low and the air is light. The rushing in of air at different temperatures from all directions produces a storm center. In order to express relationships between the volumes and weights of gases, standard conditions of temperature and pressure have been defined; these are and 76 cm. of mercury. The standard pressure adopted is the average pressure of the air at sea- level. 84. Boyle's Law. Robert Boyle, an English scientist, discovered in 1660 a law which expresses the effect on the volume of a definite amount of a gas produced by changing the pressure on it. This effect can be demonstrated by a simple experi- ment. A tube of the shape repre- sented in Fig. 10 is partly filled with mercury at a as indicated by the diagram. It is closed at b and open at c, and air is present in both arms of the tube. The pressure on the surface of the mercury at d is that of the atmosphere; since the surface at e is at the same level as FIG. 10. FIG. 10a. that at d the air entrapped in the closed end of the tube is also under a pressure of one atmosphere. The pressure on this air can be changed by pouring mercury into the tube at the point marked c. MEASUREMENT OF GASES 81 Fig. 10a represents what happens when this is done; the volume of the gas decreases. The pressure exerted on the entrapped gas can be determined by measuring the height of the column of mercury from / to g, for the following reasons : The pressure down at h is the same as that down at g. In any liquid the pressures exerted on all points at the same level are the same. The pressure down at g is equal to that of the column of mercury fg plus that of the air which presses down at /. If, in the experiment, the height of fg is 76 cm. then the pressure at g and also at h is two atmos- pheres, and we have doubled the pressure on the gas in i. If we examine the volume of the gas we shall find that it is just one-half what it was before. If a series of experiments were carried out in which the gas was subjected to pressures of 1, 2, 3, 4, 5, etc., atmospheres, it would be found that the volumes were, respect- ively, 1, i, 3, J, , etc. These facts can be stated by saying that the volume of any sample of gas varies inversely as the pressure on it; it varies inversely since as the pressure increases the vol- ume decreases. Another way to express the fact is to say that the product of the volume and the pressure is a constant, 1X1=1, J X 2 = 1, etc. These general statements can be expressed in the two forms, t> cc - or pv = c. P Where v, p, and c represent volume, pressure, and a constant, respectively. The symbol oc means " varies as." The volume varies as - because the ratio is an inverse and not a direct one. P Boyle's law is usually expressed by the equation pv = c. The law holds true only when there is no change in the temperature of the gas. 85. If we know the volume of a gas at any pressure we can calculate its volume at any other pressure or its pressure at any other volume. For any definite weight of gas pv remains constant. Let p and v represent the pressure and volume under one condi- tion, and p' and v' the values for these under another condition. Since pv = c and p'v' = c, pv = p'v'. 82 INORGANIC CHEMISTRY FOR COLLEGES If we know the values of any three quantities in this equation we can find the fourth. The solution of a problem will make this clear. A sample of gas was examined and found to have the volume . 10 c.c. when it was under a pressure of 75 cm. What would be its volume at 76 cm.? The letter p and v refer to one condition of the gas; p f and v' to the other. Let 10 = v, then 75 = p } and 76 = p r . Substitute in the formula pv = p'v' and we get 75 X 10 - 76 X x 750 = 76x x = Iff- = 9.86 c.c. Whenever a problem like this is solved it is advisable to examine the answer to determine whether it is a reasonable one, for in this way mistakes are often avoided. In this case the pres- sure is increased and, as a consequence, the volume should decrease; the answer is in accord with this and is probably cor- rect. If a mistake had been made in effecting the substitution of numbers for the letters in the formula, the answer might have been a number greater than 10, and the examination of the result would have brought out the fact that an error had been made. 86. Charles' Law. In 1787 Charles discovered that the change in volume of a gas when subjected to a change in temperature could be expressed by a simple law. When the temperature of a definite amount of a gas is changed 1 degree, the change in volume of the gas is -yfg- of the volume the gas would occupy at 0, pro- vided the pressure does not change; if the temperature is raised the gas expands, and if it is lowered it contracts. For example, if we have 1 c.c. of gas at its volume at 1 is 1 + -^3 c.c.; at 2 it is 1 + vfa c.c.; at 3 it is 1 + -^ c.c. The volume at 0, 1, 2, and 3 are f^f, f^, Iff, and f|f, respectively. These numbers are evidently in the ratio 273 : 274 : 275 : 276. If we change our scale of temperature and call centigrade 273, then 1 becomes 274, 2 becomes 275, etc. If we use such a scale of temperature we see at once that the volumes of the gas are in the same ratio as their temperatures. The temperature scale arrived at in this way was found to be of great value in other connections, and is much used in science; it is called the absolute scale. Any MEASUREMENT OF GASES 83 temperature on the centigrade scale is changed to the absolute scale by adding to it 273. A capital letter T is used to represent temperature in this scale, and a small letter I, that on the centi- grade scale; T is thus equal to t + 273. These facts can be summed up in the expression v oc T or v : v' = T : T' where v is the volume of the gas at the temperature T and v' is its volume at the temperature T'. The volumes are in the same ratios as their absolute temperatures. 87. With the aid of this law we can calculate the volume which a gas will occupy at any temperature, provided we know the volume it occupies at any other temperature. For example, if we have 20 c.c. of hydrogen at 30 what will be its volume at 40? The volume at 30 (20 c.c.) is to the volume at 40, as 30 + 273 is to 40 + 273, 20 : v' = 303 : 313 303z/ = 6260 v' = 20.6 c.c. The answer should be inspected; the temperature of the gas is increased and its volume should, therefore, be greater; the answer is a reasonable one. 88. The Combination of Boyle's and Charles' Laws. The mathematical expressions for these two laws are as follows: v oc - if the temperature is constant. P v oc T 7 , if the pressure is constant. These can be combined into one expression: v oc - X T, if pressure and temperature vary. 1 pv oc T 7 , 1 A simple analogy will make clear v/hy the volume is proportional to the product of the terms - and T. The areas of a series of quadrangles are pro- portional to their lengths if their breadths are constant; the areas are pro- portional to their breadths if their lengths are constant; and the areas are proportional to the product of their lengths and their breadths if both these factors vary. 84 INORGANIC CHEMISTRY FOR COLLEGES then pv : p'v' = T : T', or WL T_ p'v' ~ T r This formula contains six quantities; if any five are known the sixth can be determined. It is generally used to calculate the volume of a gas under certain required conditions, when the volume is known under other conditions. To do this it is changed to the following form : This means if we want to find out the volume of a gas (v) under certain conditions of temperature (T) and pressure (p) when the volume (t/) at another temperature (T'} and pressure (p') are known, we multiply this known volume by two factors, one of T_ r which is made up of the two temperatures, 7FJ , and the other of the two pressures, . This is done in a very simple way which avoids confusion in the substitution. If the temperature increases then the volume increases, and T and T' should be selected so T that the value of the fraction f is greater than 1; the larger number is put in the numerator; and the reverse is done if the temperature decreases. If, on the other hand, the pressure is to increase, the volume of the gas will be less, and the values of the pressure are substituted in the expression so that the fraction is less than 1; the smaller number is put in the numerator; if the pressure is to get less the volume will increase and the larger pres- sure should be put in the numerator. A problem will illustrate the method. What will be the volume at and 76 cm. pressure of 60 c.c. of gas measured at 20 and 74 cm. pressure? The temperatures must be expressed first in the absolute scale; they are 273 (0) and 293 (20). Since under the conditions sought the temperature is lower, the volume will be less, and T and T' MEASUREMENT OF GASES 85 are selected to make the value of , less than 1 ; therefore, we use f|f. The pressure is to increase, and the volume is, as a conse- quence, less, therefore = ff . The value of v is, accordingly, expressed thus: v = GO x m x H, 9 v = 54.4 c.c. 89. Pressure of Aqueous Vapor. Gases are often measured in vessels which are inverted over water. Under these circum- stances the vessel contains not only the gas, the volume of which is sought, but water in the form of vapor; as this is a gas it exerts a pressure and we must know the amount of this if we are to know the pressure exerted by the gas itself. That water gives off a vapor which exerts pressure in the way a gas b does, can be shown by a simple experiment. Two barometer tubes are set up side by side (Fig. 11). The level of the mercury will stand at the same height in the two tubes. A few drops of water are now forced into tube b by placing a medicine-dropper containing water under the lower end of b in the mercury. The water rises in the tube through the mercury to its upper surface, and at once the level of the mercury in it falls to c. If not too much water is added it will practically all disappear; it changes into a gas and forces the mercury down. The pressure exerted by the water- vapor is, evidently, equal to that of a column of mercury of the height equal to the distance between the level at c and at d. If next we heat the tube b we will find that the level of the mercury falls; there will be a definite place at which the mercury stands for each temperature. The difference between the height of the mercury in the barometer a and in the tube b represents in each case the pressure exerted by water-vapor at that temperature. In making such a series of observations it is necessary to have at all times a little FIG. 11. 86 INORGANIC CHEMISTRY FOR COLLEGES liquid water present, so that as much vapor as possible will form, or, as it is generally expressed, the vapor must be saturated. If the tube b is heated at 100, the boiling-point of water, the level of the mercury will be forced down to that in the dish e; this means the pressure of water- vapor is equal at its boiling-point to the pressure of the air. The pressure exerted by water-vapor, has been determined for all temperatures, and when any particular value is needed it can be found in a table. (See Appendix.) 90. If a gas is measured in a vessel over water the pressure exerted within the vessel on the surface of the water is evidently made up of the pressure of the gas plus the pressure of the water-vapor. In order to simplify the cal- culations it is advisable when measuring a gas over water to place the vessel in such a position that the level of the water inside and outside the tube is the same. In this case, Fig. 12, the pressure down at a is that of the atmosphere which is determined by reading a barometer. Since a and 6 are at the same level the pressure at 6 equals that at a and is, accordingly, equal to that of the atmosphere. The pressure at 6 is the sum of two pressures: the pressure of the gas contained in the tube and the pressure due to the water-vapor. We have then, t|Y4 =-_B- FIG. 12. pressure of the air = pressure of the gas + the pressure of water-vapor; the pressure of the gas = pressure of the air pressure of water- vapor. If the volume of a gas measured over water at a certain tem- perature and pressure is known, it is possible to calculate what the volume would be at any other temperature and pressure when the gas is free from water. This is done by using the combined expression for the two gas laws and substituting for the observed pressure of the gas, the difference between it and the pressure of water-vapor at the observed temperature. A problem will be solved as an example. A gas measured over water had the volume 50 c.c. at 20 and 740 mm. pressure. The pressure of water-vapor at 20 is 17 mm. What would its volume be if dry and at and 760 mm. pressure? MEASUREMENT OF GASES 87 EXERCISES NOTE: All temperatures given refer to the centigrade scale unless a state- ment is made to the contrary. 1. Express in degrees Fahrenheit the following temperatures: (a) 30, (6) 128, (c) 15, (d) 10. Express in degrees centigrade the following temperatures: (e) 40 F., (/) 100 F., (0) -10 F. 2. If 5 liters of air at the temperature of the room, 20, are cooled to 0, what will be the volume of the air? 3. A sample of gas was found to have the volume 250 c.c. when measured at 70 cm. pressure. What would be its volume (a) at 76 mm. pressure, and (6) if it were compressed to a pressure of 10 atmospheres? 4. The observed temperature and pressure of 50 c.c. of a gas were 23 and 745 mm. Calculate the volume of the gas at and 760 mm. 5. A sample of a gas measured over water was found to have the volume 40 c.c. when the height of the barometer was 750 mm. and the temperature was 20. Calculate the volume of the dry gas at and 760 mm. The pressure of aqueous vapor at 20 is 17.4 mm. 6. If air at 20 is passed through a hot tube and heated to 1000 what is the relation between the volume of the gas before and after passing through the tube? 7. Which is heavier, air or carbon dioxide? If you examined the air in a room lighted by gas would you expect to find a higher percentage of carbon dioxide, which is one of the products of burning gas, near the floor or the ceil- ing? Give a reason for your answer. 8. (a) Calculate the weight of 1 liter of dry air at and 760 mm. pres- sure assuming that the air contains 79 per cent of nitrogen, N 2 , 20 per cent of oxygen, C>2, and 1 per cent of argon, A. (6) What would be the volume of 1 liter if it were heated from to 20? (c) What would be the weight of 1 liter of the air at 20 and 760 mm.? 9. (a) If 1 liter of dry air at 20 and 760 mm. is shaken with enough water to saturate it with water-vapor and is then collected over water what would be the volume of the moist gas produced, neglecting the solubility of air in water? The pressure of aqueous vapor at 20 is 17.4 mm. (6) Show that dry air is heavier than moist air. 10. What is the percentage by volume of the water- vapor in the moist air prepared according to question 9 above? 11. (a) Calculate the volume of 1 gram of steam at 100 and 760 mm. What is the change in volume when 1 cram of water is changed into steam at (6) 100 and (c) at 500? 12. It was found that a certain piece of zinc when dissolved in hydro- chloric acid furnished 70 c.c. of hydrogen when measured over water at and 750 mm. Calculate the weight of zinc used. 13. What volume of hydrogen measured over water at 20 and 750 mm. pressure will be obtained when 10 grams of zinc dissolve in hydrochloric acid? CHAPTER IX WATER 91. The significance of water in nature was recognized early, for, as we have seen, the first hypothesis as to the composition of matter was that all forms of it were made up of four elements, earth, air, Ere, and water. The view that water is a constituent of matter was a reasonable one, because so many substances which are dry, and apparently free from water, yield it when heated. It was only after oxygen and hydrogen were discovered that water was shown to be a compound of these two elements. The way in which this was found out is of interest. Priestley, the discoverer of oxygen, was performing in 1781 as he said " some random experiments to entertain a few philosophical friends." He exploded some hydrogen and oxygen in a vessel by means of an electric spark. The most striking effect was that the gases decreased in volume, but it was observed that the sides of the vessel were " bedewed " with a liquid. No particular attention was paid to this fact, but Cavendish, a physicist present, was impressed by it. He considered it well worth future study, and later found that when two volumes of hydrogen were exploded with one volume of oxygen, the gases disappeared and the liquid formed was water. Lavoisier also made the same discovery. 92. Occurrence of Water. Water occurs widely distributed over the earth's surface, and as a gas in the atmosphere. In the form of ice and snow it covers lofty mountains and the surface of the earth in the neighborhood of the poles. Water is present in all living things; the fluids of the body that are essential to life- processes are chiefly water; the blood, which carries throughout the body the oxygen and other materials necessary for life, con- tains 90.3 per cent of water. The human body contains about 65 per cent of water. Without water, life as we know it would be impossible. The substances we use as foods contain a large 88 WATER 89 percentage of water; for example, there is present in potatoes 63, apples 65, string beans 90, lettuce 95, and beefsteak 62 per cent of water. The presence of such large masses of water on the earth's surface has a marked effect on the climate. When water is heated it rises more slowly in temperature than any other sub- stance, since it takes more heat to raise the temperature of water one degree than it does in the case of an equal weight of any other substance. As a consequence, islands and places near the sea have a more equable climate than those regions near the middle of a great continent. Water and ice have been the chief agencies in making the surface of the earth as it is to-day. Water slowly acts on rocks and dissolves and carries away a part of their con- stituents; some of these remain in the soil and furnish the material which plants require for their growth. Great glaciers of ice have furrowed the earth's surface, and produced valleys and moun- tain ranges. Water confined in the depths of the earth becomes heated to tremendous temperatures, and finally the enormous pressures produced disrupt great masses of the earth and as a result its surface has been greatly changed. Water is, next to coal, the most important source of energy which we can control and use for our benefit. The energy of falling water can be readily made available by allowing it to turn a wheel; the mechanical energy so produced can be utilized directly or converted into electricity, which can be distributed to distant places for use. We shall see later that one of the great recent developments in chemistry is the application of electrical energy to the preparation of many substances of value. This development has been possible only through the utilization of water-power, and more and more the value of the energy of falling water is being recognized. When all the coal is used up, in a thousand years or more, the world must turn to the waterfall for its energy, unless some new force is discovered. There are great stores of energy in the sunlight, but no one has yet discovered how to use this energy advantangeously. 93. Physical Properties of Water. Water is a liquid which has a bluish-green color when examined in thick layers. It is taken as the substance upon which to base standards for measur- ing a number of physical properties. Its freezing-point and boil- 90 INORGANIC CHEMISTRY FOR COLLEGES ing-point at 760 mm. pressure are defined as and 100, respect- ively. The weight of 1 c.c. of water at 4 was the original de- finition of a gram, but later the weight of a definite piece of platinum based on the unit, which is kept in Paris, was taken as the standard. The calorie is defined as the amount of heat re- quired to raise the temperature of 1 gram of water 1 degree, from 15 to 16; since specific heat is defined as the quantity of heat required to raise the temperature of 1 gram of a substance 1 degree, the specific heat of water is, accordingly, 1 at 15. Water is the reference substance used in expressing specific gravity (176) ; the value for the specific gravity for water is, therefore, 1. As in the case of other liquids the density (175) of water changes with temperature, but the change in this case is not uni- form; since the density of a liquid is defined as the weight of 1 c.c. of it, the density of water is 1 at 4; it decreases as the tempera- ture rises or falls from this point; at it is 0.99987 and at 100 it is 0.95838. When water freezes it expands; 1 gram of ice at has the volume 1.0908 c.c. The heat of fusion of 1 gram of ice is 79 calories, and the heat of vaporization of 1 gram of water at its boiling-point is 540 calories. 94. Water- Vapor. The fact has been brought out that when water evaporates it passes into the air as a gas. This phenomenon has a marked effect on the climate. Through the evaporation of the water on the earth's surface it passes into the air, and later, as the result of condensation, forms clouds and finally returns as rain or snow. The formation of dew when the air cools as the sun sinks and the heat reaching the earth lessens, is the result of the change of the water-vapor to a liquid. In this way growing plants get much of the water they require when it is not furnished as rain. Equally important is the fact that plants take up directly from the air water- vapor and carbon dioxide, and transform them into the organic materials of which they are made, such as wood, starch, etc. 95. The body contains much water, and that at the surface is continually evaporating and passing into the air. Our comfort is materially affected by the amount of water-vapor around us. If the air is saturated, that is, contains as much as it can hold in the form of gas, this evaporation cannot take place and we are WATER 91 uncomfortable as a result; we dread a " muggy " day. In a dry climate the water evaporates rapidly and in so doing absorbs heat from the body; as a consequence, we do not feel the effect of a high temperature as we do when the humidity is great. Air usually contains about two-thirds of the maximum amount of water it can hold as vapor; under these circumstances the relative humidity is said to be 66 per cent (two-thirds). The rate at which water evaporates is determined not by the absolute amount of water present in the air, but by the relation between the amount present and the amount the air can hold at the temperature which pre- vails. Since our comfort is determined by the rate at which water evaporates from the skin, the relative humidity is the impor- tant factor. When the relative humidity drops to 40 per cent the air feels dry, and at about 80 per cent it appears to be damp. 96. The amount of water which can exist as vapor in the air is determined by the temperature. This can be demonstrated by the use of the experiment already described in connection with the determination of the pressure of water-vapor (89). A small amount of water is allowed to rise through mercury in a barometer tube and when it reaches the surface a part evaporates and the vapor produced depresses the mercury as the result of the pressure it exerts. As the temperature is raised, more and more water evaporates and the level of the mercury sinks; if the temperature is lowered water condenses and the mercury rises. The fact that the maximum amount of water- vapor held in air varies with the temperature, is the cause of important natural phenomena. A cloud can form in a blue sky as the result of the lowering of tem- perature produced by a current of cooler air. The formation of dew has been referred to and we see now that it is the result of the drop in temperature which comes with the setting sun. In sum- mer, water soon collects on a vessel containing ice- water; this occurs when the temperature of the air next the vessel is reduced to that at which the air is saturated by the water present in it. As the temperature falls below this point the excess water precipi- tates out as drops. The temperature at which moisture is first visible on a smooth surface when air is cooled in contact with it, is said to be the dew-point of the air. 97. Chemical Conduct of Water. Water is a very stable sub- stance. When it is formed from hydrogen and oxygen a very 92 INORGANIC CHEMISTRY FOR COLLEGES large amount of chemical energy is lost as heat. In order to decompose water it must be heated to a very high temperature or brought into contact with active substances possessing a large amount of chemical energy. When water is heated to 2000 it dissociates into hydrogen and oxygen to the extent of less than 2 per cent. Water reacts with certain metals in the way already explained in section 43. In these cases the oxygen unites with the metal and hydrogen is set free; the equation for the reaction with zinc is as follows: Zn + H 2 O = ZnO + H 2 Water combines with oxides and forms bases and acids. Quick- lime, for example, is the oxide of calcium; it is converted by water into slaked lime, calcium hydroxide, which is a base: CaO + H 2 O = Ca(OH) 2 Sulphur dioxide, formed by burning sulphur in air, unites with water to form sulphurous acid: S0 2 + H 2 = H 2 S0 3 It will be recalled that it was reactions of the latter class that suggested to Lavoisier the name for oxygen; he found that when certain substances burned in the gas the products formed reacted with water to form acids; oxygen was, thus, the acid-former. It was found later, however, that many acids do not contain oxygen. 98. Hydrates. Water unites directly with certain substances and forms compounds which are more or less readily broken down, by heating at moderate temperatures, into water and the original salts. Copper sulphate, CuSC>4, is a white powder, which dis- solves in water and forms a blue solution. When the latter evaporates slowly blue crystals are formed, which have the com- position represented by the formula CuSOi, 5H 2 O. If these crys- tals are heated at a temperature somewhat above the boiling- point of water, they crumble, water is lost, and the anhydrous salt is obtained. Other salts behave in a similar way. As com- pounds which contain water separate from their solutions as crystals, the water in combination was formerly called water of WATER 93 crystallization. The formation of crystals is not dependent upon the presence of water, however. Anhydrous copper sulphate, for example, can be obtained as white, needle-like crystals. For this reason we now speak of water of hydration, rather than water of crystallization. We have anhydrous copper sulphate, and hydrated copper sulphate. To express the fact that the salt has the formula CuSC>4, 5H 2 O we call it a pentahydrate ; a salt con- taining one molecule of water of hydration is a monohydrate, one with two is a dihydrate, with three a trihydrate, etc. The pentahydrate of copper sulphate is a definite chemical compound; water is not present in it as such. It contains hydro- gen and oxygen in the proportion represented by its formula, but we do not know how the atoms are united to one another or to the other atoms present. We write the formula as we do to express the fact that when the compound is heated it readily breaks down into the compounds CuSO4 and H^O. 99. Certain crystalline substances contain water mechani- cally held within the crystal. When a piece of rock-salt is heated, water is given off. In this case the water passes into steam and sufficient pressure is produced to cause slight explosions which shatter the large crystals the substance is said to decrepitate. There is no definite relation between the weight of the water included mechanically within the crystal and the weight of the salt. With hydrated salts we can always express the relation between the weight of the water of hydration and that of the anhydrous compound by a definite chemical formula which has a quantitative significance. Most hydrated salts lose their water of hydration when heated to 100. 100. Efflorescence. If a crystal of washing soda is left in the air for some time it loses its crystalline form and changes to a white powder. This is due to the fact that at room temperature it loses a part of its water of hydration; the salt having the formula Na 2 CO3, 10H 2 O dissociates into water and the monohydrate of sodium carbonate, Na2CO,3, EbO. When hydrated copper sul- phate, commonly called blue vitriol, CuSCU, 5H2O, is left in the air, the crystals do not lose their form. We express these facts by saying that washing soda effloresces. We have learned that water gives off a vapor which passes into the air. Hydrated salts decompose so readily into water 94 INORGANIC CHEMISTRY FOR COLLEGES and the anhydrous salt that at room temperature they give off water-vapor, the pressure of which can be measured. This can be done in the same way that the pressure of water-vapor was determined (89) by placing a crystal of the substance in a barome- ter tube and noting the fall in the level of the mercury. If the pressure of the water-vapor in the air is less than the pressure of the water-vapor from the hydrated salt, the latter will decompose and water will pass into the air. Under ordinary conditions air is about two-thirds saturated with water-vapor. At 20 (68 F.) the pressure of water-vapor in the air is in the neigh- borhood of 12 mm. At this temperature the pentahydrate of copper sulphate, CuSCU, SEbO, exerts a vapor pressure of 5 mm.; it does not effloresce. The decahydrate of sodium sul- phate, Na2SO4, lOEbO, exerts a vapor pressure of 14 mm. It is seen from the above that a salt may effloresce on one day and not on another. 101. An application of hydrated salts is made in the chemical laboratory in drying gases. Anhydrous calcium chloride unites with water to form the hexahydrate CaCL?, GH^O. When air containing moisture comes in contact with the anhydrous chloride the water is absorbed and the pressure of its vapor reduced to that of the hydrated chloride; as this is exceedingly small, the air is deprived of nearly all of its water. Certain salts take up so much water from moist air that they dissolve in the water; such salts are said to be deliquescent. 102. Composition of Water. The proportion of hydrogen and oxygen in water can be determined by analysis or synthesis. In the first way the elements are separated in the free condition from water or are converted into other substances the composition of which is known; in the second, water is prepared from known amounts of other substances. The composition of water has been determined in many ways and with the greatest accuracy. A few will be sketched briefly. An experiment has been described to illustrate the fact that when hydrogen reduces copper oxide, water is formed (55). The com- position of water has been determined by a quantitative study of this reaction, which is represented by the equation, CuO+H 2 = Cu + H 2 0. WATER 95 This was done in the following way: The tube containing the copper oxide was weighed before and after the reaction took place (Fig. 13) ; the loss in weight was evidently equal to the weight of the oxygen changed into water. The amount of the water was determined by absorbing it in a tube represented at a in the diagram, which was filled with phosphorus pentoxide. This tube was weighed before and after the reaction, and the increase in weight was, evidently, the weight of the water formed. In this way the weight of the oxygen and the weight of the water were determined; the difference between these two was the weight of the hydrogen. FIG. 13. A synthesis of water has been carefully made in which all three substances involved were weighed the hydrogen, oxygen, and the water. The results of the most carefully carried out experiments lead to the conclusion that hydrogen and oxygen unite to form water in the proportion by weight of 2 of the former to 15.879 of the latter, or 2.016 to 16. 103. The composition of water has also been determined by volumetric analysis, that is, the volume relations have been studied. It has been found that 2.0027 volumes of hydrogen unite with 1 volume of oxygen to form water. If it is desired to compare the relation between the volumes of the gases which interact and the volume of the water formed, the experiment should be carried out at such a temperature that the water is a gas at 100 or above. Under these conditions it has been found that the volumes are almost exactly in the ratio of two of hydro- gen, one of oxygen, and two of steam. We shall find that this fact becomes of the greatest importance when the method of determining atomic weights is considered. 96 INORGANIC CHEMISTRY FOR COLLEGES 104. Natural Waters. Chemically pure water is very diffi- cult to obtain. Natural waters which have come into contact with the earth contain varying amounts of substances which are dissolved from the soil and the air. Rain-water is perhaps the purest natural water, but it contains dust, bacteria, ammonium compounds, and other substances which are present in the air. If it is collected after rain has fallen for some time it is practically free from these substances and is the purest water that can be obtained without special precautions; it contains, of course, the gases oxygen, nitrogen, and carbon dioxide which are present in the air. The materials dissolved from the earth find their way into the ocean, which contains about 3.6 per cent of substances which are solids. About 2.7 per cent of sea-water is common salt, the rest of the dissolved solids being chiefly chlorides and sulphates of magnesium and calcium. The Great Salt Lake of Utah contains 23 per cent of solid matter in solution. The water of lakes and rivers which are used directly as sources of supply varies in the amount of solid matter present. In places where the chief rocks are sandstone or granite, the water contains less dissolved material than that in limestone regions. We shall see later that water which contains carbon dioxide dissolves limestone slowly, and the soluble calcium salts formed render the water hard. 105. Water is freed from impurities by distillation (181). When it is heated the air dissolved in it first separates in bubbles and then escapes as the temperature is raised. At 100 it changes into steam, which is condensed by surrounding with cold water the tube through which it passes (Fig. 14). The solid materials are not volatile at 100 and remain in the flask from which the water was distilled. Water treated in this way may be considered pure enough for most purposes; it is, however, far from abso- lutely pure. To remove the last traces of foreign substances the water should be heated with a powerful oxidizing agent to destroy certain organic substances present, and be condensed in a vessel made of tin or preferably of platinum, because water dissolves traces of solid substances from glass. When left in contact with air, water slowly absorbs it. The taste of water is largely due to the air it contains; freshly distilled water tastes flat, and before being used for drinking purposes it is aerated by allowing it to flow over charcoal in the presence of air. WATER 97 106. It is often necessary to use as a water supply a river into which sewage has been deposited. If the place of pollution is far enough removed, the water may be entirely safe for domestic use, for changes take place in flowing water which result in the con- version of sewage into simple and harmless inorganic compounds. River-water usually contains suspended matter, which makes it more or less muddy and undesirable for household use. Many cities which are compelled to draw their water from rivers, subject it to an elaborate method of purification to remove suspended FIG. 14. matter and any harmful contamination produced by sewage, which often introduces disease germs into water. Cholera and typhoid fever are spread through the use of drinking water which has been infected by people suffering from these diseases. EXERCISES 1 . Why does moisture collect on a pitcher containing ice-water in summer but not in winter? 2. Calculate the weight of 1 c.c. of ice at if 1 gram has the volume 1.0908 c.c. 3. Calculate the weight in pounds of 1 cubic foot of ice. 98 INORGANIC CHEMISTRY FOR COLLEGES 4. What change in volume takes place when 100 c.c. of water at 4 are changed into ice at 0? 5. When a body floats in water the weight of the body equals the weight of the water that has the same volume as that part of the body which is under the water, (a) Taking the density of water as 1 and ice as 0.9 at 0, what would be the volume of the ice above the surface if a piece weighing 1000 grams was floated on water at 0? (6) If the block of ice is in the form of a cube, how thick would be the layer of ice above the surface of the water? (c) What percentage of the ice is above the water? (d) What weight must be placed on the ice to just cause it to sink? (e) What weight of floating ice is necessary to just support a man weighing 150 pounds? (/) If this is in the form of a cube what is its size? 6. Why is frost more likely to be formed on a cold night when the air during the day was comparatively dry than when it contained a large amount of water-vapor? 7. (a) Calculate the percentage of hydrogen and oxygen in water from the following results which were obtained in an experiment like that outlined in section 102. Weight of copper oxide used 50.250 grams, weight after hydrogen had passed over it 48.500 grams, weight of water formed 1.970 grams, (6) Calculate the percentage composition of water from its formula. CHAPTER X CHLORINE. VALENCE 107. Chlorine is an interesting and important substance; it is one of the most active elements, and enters into reactions with a great variety of other substances; it is present in a large number of compounds. A study of chlorine and the simpler substances containing it will introduce us to a number of typical chemical changes, and we shall gain thereby a deeper knowledge of how matter undergoes transformation under the influence of the chemi- cal energy which it contains and the outside energy brought to bear upon it. But there is a practical aspect also to the study of chlorine; the element and a number of its compounds find impor- tant applications which are of great service to man. It is highly probable that new uses will be found for this very active element, for activity can be made useful. 108. Occurrence of Chlorine. Chlorine is a heavy greenish- yellow gas with a stifling odor. It does not occur naturally in the free condition; for if the element were set free in the air it would soon find something with which to unite. Chlorine forms com- pounds with metals which are called chlorides; a number of these occur in nature, and some are widely distributed. Sodium chlo- ride, NaCl, which is common salt, occurs in large quantities in sea-water, of which it forms 2.7 per cent; the total amount of solids present is 3.6 per cent. Calcium chloride, CaCl2, mag- nesium chloride, MgCk, and potassium chloride, KC1, are also found in the ocean. These chlorides and other salts occur in salt deposits, which have been produced, in all probability, as the result of the drying up of inland seas. Silver chloride, AgCl, is an important source of silver. On account of its wide occurrence and its abundance, sodium chloride is the cheapest compound containing chlorine; it is, as a consequence, the substance from which chlorine is obtained on 99 100 INORGANIC CHEMISTRY FOR COLLEGES the commercial scale, either directly or indirectly. This fact is expressed briefly by saying that sodium chloride is the source of chlorine. 109. History of Chlorine. Scheele, a Swedish apothecary, first separated chlorine from one of its compounds in 1774. He was studying a mineral called pyrolusite and found that when it was treated with muriatic acid, a yellow gas with a stifling odor was formed. Pyrolusite was shown later to be a dioxide of an element called manganese, and to have the formula Mn02. Muri- atic acid was so called because it was obtained from salt, which is found in the sea (Latin murias)', it is what is now called hydro- chloric acid and has the formula HC1. The new substance was named chlorine (from x^P^, greenish-yellow). On account of the fact that it was prepared with the aid of a substance containing oxygen it was thought to contain this element, but after a very thorough study of chlorine by Sir Humphry Davy it was recog- nized as an element. 1 110. Preparation of Chlorine* (a) Electrolysis of Hydro- chloric Acid. Chlorine cannot be conveniently prepared by heat- ing a chloride; mercuric oxide yields oxygen when heated, but chlorine cannot be prepared by a similar decomposition of mer- curic chloride. We accordingly use another form of energy to effect the decomposition. When electricity is passed through a solution of hydrochloric acid, which is a compound of hydrogen and chlorine, the acid is decomposed into its constituents. The experiment can be carried out in a Hofmann apparatus, but the platinum electrodes should be replaced by ones of carbon, since chlorine unites directly with the metal to form platinum chloride, which is soluble in water, and, as a consequence, the valuable metal is slowly dissolved. When the experiment is carried out and precautions are taken to eliminate a disturbing factor intro- duced through the fact that chlorine is much more soluble in water than hydrogen, it will be found that as a result of the decom- position the two gases are formed in equal volumes. This method of making chlorine recalls one of the methods of preparing oxygen from water. In either case the compound of the element with hydrogen was decomposed into its constituents by electricity. 1 Davy's description of his experiments can be found in Volume 9 of the Alembic Club Reprints. It is of interest to read by what experimental and logical processes a substance was shown to be an element. CHLORINE. VALENCE 101 Chlorine is prepared on the commercial scale by the electrolysis of solutions of sodium chloride; as a result, in addition to chlorine are obtained sodium hydroxide and hydrogen, both of which are valuable products. The details of fe : process ave j , given in sec- tion 605. 111. (6) Deacon Process. Chjo#n$ enrobe' ser?a1^t6d.from the hydrogen with which it is united in hydrochloric acid, through the agency of chemical energy. When a mixture of air and hydro- chloric acid heated to about 400 is passed through a tube filled with clay balls which have been soaked in a solution of copper chloride and then dried, chlorine is formed as the result of a reac- tion expressed by the following equation: 4HC1 + O 2 = 2H 2 O + 2C1 2 The finely divided copper chloride deposited on the balls of clay serves as a catalytic agent. Chlorine is prepared in this way by the oxidation of hydrochloric acid; oxygen is the oxidizing agent, and the products of oxidation are water and chlorine. Chlorine made by this process is mixed with the nitrogen gas present in the air which furnished the oxygen; the amount of nitrogen is great, since four-fifths of any volume of air is this gas. This way of making chlorine is known as the Deacon process; it was formerly much used in England to prepare chlorine for making bleaching powder, but other methods yield chlorine in a purer condition, and are now used almost exclusively. 112. (c) Manganese Dioxide and Hydrochloric Acid. It has been stated that no chlorides are available which yield chlorine readily on heating. Chlorides can be prepared, however, which are so unstable that they undergo decomposition spontaneously when they are formed. Reactions of this type furnish conven- ient methods of preparing chlorine. One of these reactions is the one which led to the discovery of the gas. When manganese dioxide is treated with hydrochloric acid, chlorine, manganese chloride, and water are formed. It is believed that the reaction takes place in two stages; manganese tetrachloride is first formed, and then breaks down into manganese dichloride and chlorine: MnO 2 + 4HC1 = MnCU + 2H 2 O MnCU = MnCl 2 + C1 2 102 INORGANIC CHEMISTRY FOR COLLEGES If the reaction is carried out at a very low temperature manganese tetrachloride is formed; if now, the temperature is allowed to rise the tetrachloride loses chlorine in the way indicated by the last equation ab>ve.\ A single;; equation can be written to express the two reactions. Sinc'e the 'man'ganese tetrachloride decomposes as soon as formed,; ife-fbrmijlja should nbt appear in the equation which is to represent trie 'formulas o'f the substances used and those obtained. The equation is, accordingly, MnO 2 + 4HC1 = MnCl 2 + C1 2 + 2H 2 O Certain other dioxides, such as lead dioxide, behave in a similar way when treated with hydrochloric acid. 113. Large quantities of chlorine are manufactured in Eng- land by the reactions indicated by the equation given above. Manganese dioxide occurs abundantly in nature and is conse- quently cheap in price, and hydrochloric acid is a by-product which is produced in large quantities in the industrial method of manufacturing soda used in England. Soda is not made in this way in America, and, as a consequence, hydrochloric acid is not available at a low price. Water-power to make electricity is available, however, and, as a result, chlorine is manufactured in America almost exclusively by the electrolysis of sodium chloride. This process yields a valuable by-product, sodium hydroxide, NaOH, which is commonly called caustic soda in trade. It is evident that the particular processes used to prepare chemical substances on the large scale for commercial purposes are deter- mined in the last analysis by cost, in which must be taken into account such factors as source of materials, expense of transporta- tion and labor, availability of energy derived from heat and elec- tricity, value of by-products produced, etc. The preparation of useful substances from this economic point of view is considered in industrial chemistry. 114. Chlorine is prepared in the laboratory from hydrochloric acid and manganese dioxide, or by a slight modification of this method. Instead of using hydrochloric acid which has been pre- pared previously, this compound is made in a vessel containing manganese dioxide. The preparation of hydrochloric acid will be discussed in detail later; it is sufficient to note here that the com- pound is formed when sodium chloride, which is common salt, is CHLORINE. VALENCE 103 treated with sulphuric acid. The reaction that takes place, in which sodium sulphate and hydrochloric acid are formed, is repre- sented by the following equation: 2NaCl + H 2 SO 4 = 2HC1 + Na 2 SO 4 It should be observed that the sodium and hydrogen atoms change places. The hydrochloric acid produced in this way reacts with the manganese dioxide present according to the equation already given : MnO 2 + 4HC1 = MnCl 2 + C1 2 + 2H 2 O This reaction leads to the formation of manganese chloride. Since all chlorides react with sulphuric acid and are converted into sulphates, the reaction indicated below takes place: MnCl 2 + H 2 SO 4 = 2HC1 + MnSO 4 The three equations given above represent what occurs when manganese dioxide, sodium chloride, and sulphuric acid are mixed. The products obtained are sodium sulphate, Na 2 SO 4 , manganese sulphate, MnSO 4 , water, and chlorine. Hydrochloric acid and manganese chloride, MnCl 2 , are so-called intermediate products; they react with other substances as soon as formed a fact indi- cated by the appearance of their formulas on the left-hand side of the equations; they are not final products and their formulas do not appear in the final equation. A single equation can be written to express the sum of all these reactions. First the formulas of the substances used are set down, then an arrow, and next the formulas of the final prod- ucts of the reaction, thus : NaCl + H 2 SO 4 + MnO 2 -> Na 2 SO 4 + MnSO 4 + H 2 O + C1 2 The expression is next balanced beginning with the element sodium, Na. This is done by putting a 2 in front of the formula NaCl, since two sodium atoms are present in sodium sulphate, Na 2 SO 4 . Having balanced the sodium in sodium sulphate we next balance the group of atoms, SO 4 , which is present in this compound, and also in manganese sulphate and in sulphuric acid; this is done by taking two molecules of sulphuric acid ; the expres- sion becomes as the result of these changes, 2NaCl + 2H 2 SO 4 + MnO 2 -> Na 2 SO 4 + MnSO 4 + H 2 O + C1 2 104 INORGANIC CHEMISTRY FOR COLLEGES Proceeding with the acid in which the SO 4 group has been bal- anced, we next examine the number of hydrogen atoms; these are balanced by putting 2 in front of the formula of water. The other atom present in this compound, oxygen, is next balanced; there are now two atoms on the right of the arrow and two on the left in combination with manganese. As no change is necessary at this step we balance the manganese; no change is required here. The remaining atoms, chlorine, are examined and found also to be balanced. The expression becomes as the result of all these changes, 2NaCl + 2H 2 SO 4 + MnO 2 = Na 2 SO 4 + MnSO 4 + 2H 2 O + C1 2 The arrow has been replaced by the sign of equality, since a re-examination of the entire equation shows that all atoms are balanced. The student should carefully study the process by which the equation was balanced. It. is of the greatest importance that facility should be gained in carrying out the process, for when this is acquired, the writing of chemical equations offers little difficulty. 115. There are a number of substances besides manganese dioxide which react with hydrochloric acid and produce chlorine. One of these, potassium permanganate, can be used conveniently for this purpose; if concentrated hydrochloric acid is allowed to fall on it, drop by drop, at room temperature, the gas is formed. The cost of potassium permanganate interferes with the general use of this way of making chlorine. 116. (d) Bleaching Powder and Acids. Chlorine can be pre- pared conveniently by treating bleaching powder with an acid. Bleaching powder is made by treating slaked lime with chlorine. The formula of lime, calcium oxide, is CaO, and that of bleaching powder CaOCl 2 ; for this reason it is often called chloride of lime; it has important uses which will be described later. The reactions which take place between this substance and hydrochloric acid and sulphuric acid, respectively, are represented by the following equations : CaOCl 2 + 2HC1 = CaCl 2 + H 2 O + C1 2 CaOCl 2 + H 2 S0 4 = CaS0 4 + H 2 O + C1 2 CHLORINE. VALENCE 105 In the first case calcium chloride is formed, along with water and chlorine: in the second, calcium sulphate. It is to be noted that all substances which contain the group of atoms represented by SO4 are called sulphates; applying this statement to sulphuric acid, H2SO4, it should be called hydrogen sulphate a name which is frequently used for the acid. Chloride of lime is readily available, as it can be purchased at grocery stores. For this reason it is a convenient source of chlo- rine if it is desired to use it in the household for bleaching or other purposes. Vinegar contains an acid, acetic acid, and can be used to liberate the chlorine. 117. Properties of Chlorine. Chlorine is a heavy, greenish- yellow gas which has a stifling odor. It affects seriously the lining of the throat, nose, and lungs, and when inhaled even in moderate quantities produces the effects which result from a bad cold. In larger amounts it produces suffocation and death. On account of the3e effects chlorine was used in the recent war to disable an army before an attack was made. The gas can be liquefied, and when stored in cylinders of iron is readily transported. When the valve of a cylinder containing liquid chlorine is opened the substance rushes out in the gaseous condition. As it is a very heavy gas it stays near the surface of the earth and when impelled by the wind it rolls on as a deep, yellowish-green cloud. Certain compounds prepared from chlorine and other substances affect the eyes, causing intense pain and a copious flow of tears; such sub- stances were used in warfare to help rout the enemy. In order to withstand gas-attacks all soldiers were furnished with rubber gas-masks, which fitted the face closely and were connected with a cannister that contained substances to absorb the gases. Espec- ially prepared charcoal which absorbs into its pores large quanti- ties of gas was used for this purpose, together with a mixture con- taining lime and sodium hydroxide which reacts with chlorine and the other gases containing this element. 118. One liter of chlorine at and 760 mm. pressure weighs 3.220 grams. It is 2.49 times as heavy as air. Liquid chlorine has the specific gravity 1.41 at 20; it boils at 33.6 and changes to a solid at 102. The pressure of liquid chlorine when stored in cylinders is 6.6 atmospheres at 20. When chlo- rine is passed into water at 20, 215 volumes of the gas dissolve 106 INORGANIC CHEMISTRY FOR COLLEGES in 100 volumes of the liquid; the solubility at is 461 volumes in 100; for this reason the gas is not collected in the laboratory over water, in the way used in the case of hydrogen and oxygen. It is collected by the upward displacement of air; when chlorine is conducted into a bottle filled with air, the heavy gas settles to the bottom and as it rises forces the air out. 119. Chemical Conduct of Chlorine. With Phosphorus. The chemical conduct of chlorine can be best illustrated by a number of striking experiments. Several jars of the gas are provided. Into one is introduced a bit of phosphorus; it soon reacts with the production of light. On account of this fact we say that phos- phorus burns in chlorine. The definition of the word burn has been broadened to include all chemical phenomena which take place with the evolution of light and heat; it is not now restricted to a change involving oxygen. When phosphorus burns in an atmosphere of chlorine, phosphorus pentachloride, a yellow solid, is formed : 2P + 5C1 2 = 2PC1* It will be recalled that it yielded phosphorus pentoxide when burned in oxygen. With Antimony. If powdered antimony is dropped into a jar containing chlorine the two elements react and a brilliant flash is seen where each particle of the powder unites with the chlorine. Antimony trichloride is formed: 2Sb + 3C1 2 = 2SbCl 3 When antimony is burned in oxygen, antimony trioxide, Sb 2 O3, is the result of the reaction. With Copper. When a piece of copper in the form of a thin foil is gently heated and then introduced into chlorine, union takes place with the evolution of light, and copper chloride is formed : Cu + C1 2 = CuCl 2 When copper is heated in oxygen, copper oxide, CuO, is produced. With Sodium. Sodium, as we have seen, is a very active metal which decomposes water with great violence; chlorine is a poisonous gas. When these substances unite we get as the result sodium chloride, common table-salt. The experiment is an easy CHLORINE. VALENCE 107 one to carry out. Some of the metal is rolled into a thin sheet, warmed a little, and put into a jar of chlorine; it burns and a white powder is formed which is recognized by its taste. The equation for the reaction is, 2Na + C1 2 = 2NaCl 120. With Hydrogen. When a jet of burning hydrogen is put into a jar of chlorine it continues to burn, this time, however, with a luminous flame, until the chlorine is used up a fact that can be recognized by the disappearance of the yellow color. The product is a colorless gas, hydrogen chloride, which, when dissolved in water, is called hydrochloric acid: H 2 + C1 2 = 2HC1 The fact that an acid is formed can be shown by testing the gas with a moist piece of blue litmus paper, which changes to pink in the presence of acids. Hydrogen and chlorine can be made to unite through the action of light. If a thin- walled glass vessel is filled with a mix- ture of these gases and placed in the direct sunlight, the union takes place with explosive violence. Chlorine and many of its compounds are susceptible 'to the action of light; we shall see later that a number of photographic processes are based upon this fact. When a flame is introduced into a mixture of hydrogen and chlorine the gases unite with the production of a loud noise. The behavior of hydrogen with chlorine recalls in many ways the action of this element with oxygen. 121. With Other Elements. Chlorine reacts with nearly all the elements, chlorides being formed more or less rapidly at the ordinary temperatures. Carbon resists the action of chlorine and is employed in making electrodes to be used in the preparation of the gas by electrolytic methods. If chlorine and the element with which it is to react are both carefully dried before being brought together, the action takes place very slowly in many cases. Thus, bright metallic sodium can be kept without change in an atmos- phere of dry chlorine for a long time. Water in this case is a catalyst, which markedly affects the rate at which the chemical change takes place. Water often acts in this way, for many sub- stances which react rapidly,- and sometimes with violence, under 108 INORGANIC CHEMISTRY FOR COLLEGES ordinary conditions, do not appear to affect each other if water is rigorously excluded. 122. With Hydrocarbons. The experiments with hydrogen and chlorine which have been described show that these elements possess a great affinity for each other the chemical energy bound up with each element is of such a nature that it tends to escape as heat when the opportunity is offered. Not only does chlorine react with free hydrogen, but it withdraws hydrogen from com- pounds containing it. Hydrocarbons belong to a class of sub- stances which are composed of hydrogen and carbon a fact which gives them their name. Turpentine is a familiar hydro- carbon which is obtained from pine trees. If some of this sub- stance is warmed and poured on a piece of paper which is then put into a jar containing chlorine, a striking phenomenon takes place; a dense black cloud of very finely divided carbon is formed; and a test of the contents of the jar will show that hydrogen chloride is present. The reaction consists in the union of the chlo- rine and hydrogen, and the carbon is left behind. Other hydro- carbons and compounds containing carbon, hydrogen, and oxygen, react in a similar way. When a burning candle is lowered into a jar of chlorine, carbon is deposited as it burns, and hydrogen chloride is formed. Other compounds are produced, however, and the phenomenon is not so striking as when turpentine is used. When most hydrocarbons and compounds which contain carbon, hydrogen, and oxygen are subjected to the action of chlorine, hydrogen chloride is formed, and chlorine enters the molecule and takes the place of the displaced hydrogen. For example, when methane, CH4, which is the chief constituent of natural gas, is treated with chlorine such a reaction takes place: CH 4 + 4C1 2 = CC1 4 + 4HC1 This particular kind of reaction is called substitution, since one element takes the place of another. Reactions of this type are greatly facilitated by sunlight. The discovery that one element could take the place of another in this way was the result of an unusual occurrence. A state ball was being held at the Tuilleries in the reign of Louis Philippe. Soon after the candles were lighted the room was filled with suffocating fames. The guests were dis- missed, and the cause of the remarkable occurrence was investi- CHLORINE. VALENCE 109 gated. The master of ceremonies happened to be a relative of a young chemist, Dumas by name, who was asked to find out the source of the fumes. He found that the candles had been bleached by chlorine, and that some of the element had driven out a part of the hydrogen in the wax of which the candles were made, and had taken its place. When the candles were afterwards burned this chlorine was given off in combination with hydrogen, as hydrogen chloride; and this was the substance which caused the discom- fiture of the king and his guests. This type of reaction was studied carefully by Dumas, and became of the greatest importance in organic chemistry. 123. With Water. If water through which chlorine is passing is cooled to 0, a white, crystalline substance separates. It is called chlorine hydrate, and has the formula Ck, 8H20; the comma placed between the formula of chlorine and that of water is used to indicate the fact that the compound readily decomposes into the constituents of which it is made up. Chlorine hydrate belongs to the class of substances often called molecular com- pounds, i.e., compounds formed as the direct union of two or more molecules, which more or less readily break down into these same molecules. Water forms many such molecular compounds, which are commonly called hydrates. Sulphuric acid, for example, forms a hydrate for which the formula is EkSCU, H^O. 124. Michael Faraday, who worked in the Royal Institution in London, a place where many great discoveries in physics and chemistry have been made, was studying (1823) chlorine hydrate when he hit upon a method of liquefying chlorine. He heated some of the hydrate in a closed tube shaped like an inverted letter V; one arm of the tube contained the hydrate and the other was surrounded by water. A yellow oil which appeared in the cold part of the tube proved to be pure chlorine in the form of a liquid. It was produced as the result of the fact that when the chlorine hydrate was heated it dissociated into its constituents, and the gas liberated in such a small volume created a great pres- sure and, as a consequence, turned to a liquid. This observation led Faraday to study other gases, and he finally succeeded in liquefying a number by subjecting them to pressure at low temp- eratures. Many years after, Dewar took up the work again in the Royal Institution and with the aid of improved methods and 110 INORGANIC CHEMISTRY FOR COLLEGES the knowledge gained as chemistry and physics advanced, liquefied the gases which did not yield to Faraday's attempts. 125. A solution of chlorine in water contains, in addition to chlorine, some hydrochloric acid and a compound of the com- position represented by the formula HOC1, called hypochlorous acid. The substances are formed as the result of the reaction represented as follows : H 2 O + C1 2 = HOC1 -4- HC1 In this case one-half the hydrogen in the molecule of water is substituted by chlorine, HOH > HOC1. All the chlorine does not disappear from the solution, however, because hypochlorous acid and hydrochloric acid interact to form water and chlorine, the equation for the reaction being the reverse of that just given: HOC1 + HC1 = H 2 O + C1 2 We have in this case what is called a reversible reaction one in which certain substances interact to form others which, in turn, interact to form the original substances. This is expressed briefly by replacing the sign of equality in the equation by the symbol ^, thus: If water and chlorine are brought together they interact to form hypochlorous acid and hydrochloric acid, and these in turn to form chlorine and water; the solution contains, as a consequence, all four substances. When one volume of chlorine is dissolved in one volume of water at 10 about 66 per cent of the chlorine is present as free chlorine and the rest as hypochlorous acid and hydrochloric acid. This fact can be represented thus: 34 H 2 O + C1 2 <= HOC1 + HC1 66 If a solution of chlorine in water (chlorine-water) is exposed to direct sunlight the hypochlorous acid present decomposes according to the following equation: 2HOC1 = 2HC1 -f 2 CHLORINE. VALENCE 111 As the oxygen is slowly set free it separates from the solution in bubbles and can be collected in a suitable apparatus. As soon as the hypochlorous acid begins to decompose more chlorine reacts with the water present to form the acid; this process continues until all the chlorine disappears, provided the reaction takes place in sunlight; the complete change which takes place is represented the following equation: 2H 2 O + 2C1 2 = 4HC1 + 2 This reaction is the reverse of that written for the Deacon process for the manufacture of chlorine; it is, thus, a reversible reaction, and it is for this reason that when oxygen and hydrogen chloride react they do not change completely to chlorine and water a fact which has been noted. The reaction which takes place in the Deacon process is brought about when the substances are in the gaseous condition; the reaction in the reverse direction proceeds best when chlorine is dissolved in water. The conditions under which a reversible reaction takes place influence greatly the extent to which it proceeds in either direction. This important fact will be discussed later. 126. Chlorine as an Oxidizing Agent. When substances are treated with chlorine and water they are oxidized, the oxygen for the purpose being furnished by the hypochlorous acid present in the solution. We have not as yet become familiar with such substances, but a simple case can be cited as an illustration. Sulphurous acid has the formula H^SOs; it can be changed by oxygen oxidized to sulphuric acid, H 2 SO4. The changes which occur when sulphurous acid is treated with chlorine and water can be expressed by the following equations: C1 2 + H 2 O = HOC1 + HC1 HOC1 = HC1 + O H 2 SO 3 + O=H 2 SO 4 These equations can be combined into one by the process which has been described at length (71) : C1 2 + H 2 O + H 2 SO 3 = 2HC1 + H 2 S0 4 112 INORGANIC CHEMISTRY FOR COLLEGES 127. It should be noted that we do not write the second equation above in the form used before : 2HOC1 = 2HC1 + O 2 The formula for oxygen gas is 02. When chlorine and water act as an oxidizing agent no oxygen gas is given off; we assume, therefore, that as soon as an oxygen atom is set free it immediately unites with the sulphurous acid and forms sulphuric acid according to the third equation above. There is experimental evidence in favor of this view. When oxygen is formed in the presence of a substance that can be oxidized, it is more reactive than oxygen gas. Many substances which do not react with the latter at ordinary temperatures, do react with oxygen if the element is liberated from a compound in the presence of these substances. In order to make this distinction, we say that the element when liberated from a compound is in the nascent state, the word nas- cent being derived from the Latin word nascens, born. Chlorine and water furnish, thus, a convenient source of nascent oxygen. The explanation of these facts from the theory of atoms is that the oxygen when set free is in the form of atoms the symbol for nas- cent oxygen is O ; if nothing is present to combine with these atoms they unite and form molecules, and thus produce oxygen gas, the formula for which is 62 . The activity of nascent oxygen is explained also from the standpoint of energy. When the atoms are set free they possess a large amount of chemical energy and are, therefore, active. If nothing is present to be oxidized these atoms lose a part of their energy when they react with each other to form molecules. This reaction can be expressed thus: 2O = O 2 When it takes place a large amount of chemical energy is lost and the resulting molecule is. accordingly, less active. 128. Similar relationships are observed in the case of hydrogen. Many substances that are not affected by hydrogen gas are re- duced when placed in a vessel in which hydrogen is being formed as the result of the action of an acid on a metal. A similar explana- tion is offered. Atoms of hydrogen are first liberated; these are CHLORINE. VALENCE 113 active; if no reducible substance is present they combine, lose a part of their chemical energy, and the resulting molecule is less active than the atoms of which it is composed. What occurs according to this view when zinc and hydrochloric acid react can be expressed by equations as follows: Zn + 2HC1 = ZnCl 2 + 2H 2H = H 2 The hypothesis put forward to account for the nascent state is valuable since it offers an explanation of facts other than those which led to its suggestion. It will be recalled that gaseous hydro- gen does not reduce copper oxide at room temperature to copper and water; the substances must be heated together. It is highly probable that at the higher temperature the molecules of hydrogen break down in part into atoms; these, according to the hypothesis, are more active than molecules, and, as a consequence, they with- draw the oxygen from the copper oxide. If this view is correct the molecules of hydrogen first dissociate into atoms; the reaction which takes place can be represented as follows; H 2 = 2H It has been shown that many molecules can be decomposed by heat into parts which reunite when the source of heat is removed; when this decomposition takes place the molecules are said to dissociate. Water, for example, dissociates to the extent of 1.8 per cent at 2000 into hydrogen and oxygen. If the mixture of the three substances at this temperature is cooled, the hydrogen and oxygen recombine to form water. 129. Chlorine as a Bleaching Agent. If pieces of dyed cotton cloth which have been previously dried are suspended in a jar of dry chlorine, scarcely any change takes place. If other pieces of the same cloth are dipped in water and then subjected to the action of chlorine in the same way, a change is soon observable; the dyes begin to fade, and after a few minutes the pieces of cloth become white or nearly so. These experiments show that chlorine and water destroy certain dyes. They affect in a similar way the coloring matter present in unbleached cotton, and could be used 114 INORGANIC CHEMISTRY FOR COLLEGES for bleaching. The reaction which takes place is expressed, in the main, by the equations already given: H 2 O + C1 2 = HOC1 + HC1 HOC1 = HC1 + O The nascent oxygen then oxidizes the colored compound, which, as a result, is converted into other substances that are colorless. 130. Test for Chlorine. When free chlorine is brought into contact with a solution of potassium iodide, KI, the iodine is set free and potassium chloride is formed: 2KI + C1 2 = 2KC1 + I 2 The iodine can be recognized even when it is present in very small amounts by shaking the solution with carbon disulphide; it dis- solves in the latter and imparts to it a characteristic purple color. The iodine can also be detected by treating it with a solu- tion made by heating a little starch with water; in this case a blue color is produced. These reactions can be used to detect chlorine, but since other substances liberate iodine from potassium iodide they alone do not prove the presence of the free element. In order to draw a definite conclusion the solution is tested for the chloride formed, by the method described in section 145. 131. Uses of Chlorine. Most of the chlorine that is manu- factured is converted into chloride of lime, CaOCk, or sodium hypochlorite, NaOCl, compounds which are extensively used in bleaching. Chloride of lime yields chlorine readily and is used as a disinfectant, a deodorizer, in the manufacture of chloroform, and for other purposes. Free chlorine in water destroys bacteria; the gas is, accordingly, now used in sterilizing water supplies when contamination is suspected or has been shown to be present. Chlorine is used for making chlorides of many of the elements, in the preparation of potassium chlorate, and in separating other elements from their compounds, such as bromine and iodine. Tin is recovered from waste tin cans by an ingenious method which involves the use of chlorine. Tin cans are made from sheet-iron which has been covered with tin. As the latter is a more or less expensive metal it is desirable to recover it from used cans. The problem was finally solved in a simple way. When CHLORINE. VALENCE 115 chlorine was passed over the cans, the tin was converted into the chloride, SnCU, which boils at 114. The mixture was then heated and the chloride distilled off. The product, stannic chloride, is used in large quantities as a mordant in dyeing silk. Chlorine is also used in the preparation of carbon tetrachloride, a substance which has recently been shown to be very useful on account of the fact that it is a liquid that readily passes into a vapor which is non-inflammable and, the presence of which pre- vents combustible substances from burning. It is used in fire extinguishers, and when added to gasoline makes a mixture that can be used for cleaning purposes without the danger which attends the use of gasoline alone. 132. Comparison of the Chemical Conduct of Chlorine with that of Oxygen. The acquisition of the large number of facts that one meets in the study of chemistry is greatly facilitated by exam- ining new facts as they are presented in the light of the knowledge already gained. This is best done by searching out resemblances and differences, and by determining whether the new facts are additional examples of general principles that have been learned. This method will be illustrated by an examination of the chem- istry of oxygen and chlorine. First resemblances will be noted, then differences. Occurrence: Both elements are active and occur in combina- tion with other elements; oxides and chlorides of metals are important substances found in nature. A much larger proportion of oxygen than of chlorine is present in the earth; oxygen is found in the free condition, chlorine is not. Preparation: Both elements can be obtained by heating the compounds which they form with inactive metals; platinum oxide yields oxygen and platinum chloride yields chlorine when heated. If the metal in combination with chlorine or oxygen is an active one, such a decomposition does not take place; we cannot decom- pose sodium chloride or sodium oxide by heat. Some chlorides decompose more readily than the corresponding oxides, and some do not; manganese tetrachloride breaks down spontaneously into manganese dichloride and chlorine; in order to decompose man- ganese dioxide a high temperature is required. On the other hand, mercuric oxide is decomposed into mercury and oxygen at a lower temperature than that at which an analogous decomposi- 116 INORGANIC CHEMISTRY FOR COLLEGES tion takes place in the case of mercuric chloride. Both elements can be prepared by the action of an electric current on their com- pounds with hydrogen; hydrogen oxide, water, under these cir- cumstances yields hydrogen and oxygen, and hydrogen chloride yields hydrogen and chlorine. Chemical conduct : The two elements unite with most other ele- ments and form oxides and chlorides. The phenomena observed in the two cases resemble each other very closely; hydrogen burns in oxygen, also in chlorine; phosphorus and other elements behave in a similar way; but carbon burns in oxygen and not in chlorine. The kindling temperature for a reaction between an element and chlorine is usually lower than that of the analogous reaction with oxygen. Most substances react with chlorine at room tem- perature more rapidly than with oxygen. In general, chlorine is the more active element at ordinary temperatures. Chlorine withdraws hydrogen from compounds containing it, and unites with the hydrogen and also with the element to which the hydro- gen was joined. Oxygen behaves in a similar way. Equations for two typical reactions are as follows : CH 4 + 4C1 2 = CC1 4 + 4HC1 CH 4 + 2O 2 = CO 2 + 2H 2 O It will be seen from the comparisons which have been drawn that chlorine and oxygen resemble each other very closely in chemical properties. The facts which have been given in the discussion of the two elements can be remembered more readily and appreciated more fully as the result of this comparison. At this stage in his study of chemistry the student would probably not be able to make as full a comparison as that given above, but as he advances he will soon learn to generalize his facts if he con- stantly examines carefully new facts as they are presented. He should get in the habit of asking, is this like anything I have learned before, can I hang this new fact on an old one? Of course this most efficient method of study can be used only when the subject is studied from day to day; if the facts are not acquired as they are presented, they cannot evidently be used to help remember new facts. A real, usable knowledge of chemistry can be gained only through continuous study; in no other way can one get the benefit that is won through the study of a science. CHLORINE. VALENCE 117 133. Valence. A comparison of the formulas of a series of compounds containing chlorine with those of a series of com- pounds of the same elements with oxygen, brings out some striking relations. Let us examine a few of these set side by side: HC1 H 2 O NaCl Na 2 ZnCl 2 ZnO CaCl 2 CaO A1C1 3 A1 2 3 SbCl 3 Sb 2 O 3 MnCU Mn0 2 ecu CO 2 We see from the first column that 1 hydrogen atom unites with 1 chlorine atom, and that 1 sodium atom unites with 1 chlorine atom; zinc and calcium each unite with 2 chlorine atoms; alumi- nium and antimony with 3; and manganese and carbon with 4. The elements differ in the number of atoms of chlorine which they can hold in combination. An examination of the second column shows not only that these same elements differ in their power to hold oxygen atoms, but that there is remarkable relationship be- tween the combining power of any element as measured by chlor- ine in one case and by oxygen in the other. This will be clear from the following considerations : Hydrogen is taken as the stand- ard of combining power when we wish to consider the number of atoms which unite with one another. We say the combining power of hydrogen is 1 or use a more technical word and say its valence is 1, the word being derived from the Latin word valentia signifying strength. Since 1 chlorine atom unites with 1 hydrogen atom its valence is also 1. We, thus, have two unit standards by which we can measure the valence of any other element. We can now readily express the differences in combining power observed among the elements which occur in the formulas listed in the first column above. Hydrogen and sodium each have the valence 1, zinc and calcium have the valence 2, aluminium and antimony have the valence 3, and manganese and carbon the valence 4. These conclusions are drawn from the study of the for- mulas of the chlorides of these elements. We shall next examine the oxides. One oxygen atom unites with 2 hydrogen atoms; it, 118 INORGANIC CHEMISTRY FOR COLLEGES accordingly, has the valence 2. Knowing this we can determine the valence of the elements in the compounds in the second column by comparing them with oxygen and remembering that its valence is 2. In the formula Na 2 O, 2 sodium atoms are represented in combination with 1 oxygen atom ; the combining power of the lat- ter which is 2 is used up in holding 2 sodium atoms; each of the lat- ter must have a combining power of 1 the valence of sodium is 1 . In zinc oxide, ZnO, the valence of zinc is 2 because oxygen has this valence, and 1 atom is in union with 1 atom. It must not be concluded that the valence of zinc is 1 because it unites with 1 oxygen atom. We must always go back to the standard selected, namely hydrogen, and in this standard the combining power of oxygen is 2, as has been stated; therefore, that of zinc is, also 2. Calcium is like zinc in valence. When we come to aluminium oxide the case appears to be more difficult. It is simple if looked at in the following way: As each oxygen atom has the combining power of 2, 3 atoms of this element have the combining power of 6; this is used in holding in combina- tion 2 atoms of aluminium; and each atom of this element must have, as a consequence, the valence 3; likewise antimony has the valence 3. Manganese and carbon each unite with 2 oxygen atoms and, consequently, they have the valence 4. As a result of the study of the formulas of the chlorides of a series of elements, we come to a conclusion as to the combining power of these elements; a study of the formulas of the oxides leads to conclusions as to their combining capacity identical with those arrived at in the first way. This is a striking fact, and is more remarkable when we find that the valence which each of these elements possesses in other compounds is the same as that arrived at from a consideration of the formulas of their chlorides. The valence of an element is thus an expression of a chemical characteristic of that element. We shall see that this fact sim- plifies markedly remembering chemical formulas; if we know the valence of the elements which are present in a compound we can, in most cases, write its formula. A few examples will make this clear. Zinc always has the valence 2. What will be the for- mulas of the compounds of zinc with bromine, sulphur, and phos- phorus? The symbols for these elements and the valence of each i ii in are represented thus Br, S, P. A Roman numeral is often placed CHLORINE. VALENCE 119 over a symbol to represent the valence of the element. When the writing of chemical formulas becomes " second nature " these numbers can be omitted; but until the student is sure of himself, he should make it a practice to indicate the valence of the elements, for such a practice will prevent many mistakes. The formulas of the compounds of the elements with zinc are evidently as follows: ii i ii ii ii in ZnBr 2 ZnS Zn 3 P 2 The sum of the combining capacities of all of the atoms of one element in the compound must equal the sum of the combining capacities of all of the atoms of the other. In the case of zinc bromide, 1 zinc atom with the valence 2, 1 X 2 = 2, is united with 2 atoms of bromine each having the valence 1, 2X1=2; in the second case, zinc, 1X2=2, and sulphur, 1X2=2; in the third compound there are 3 atoms of zinc, 3X2=6, and 2 atoms of phosphorus, 2X3=6. As a mathematical check on the correctness of a formula we can multiply the number of atoms of one element present by its valence, and the number of atoms of the second element by its valence, and we must get the same number in each case. It is evident that the following do ii i ii ii ii ii not represent chemical compounds: ZnCls, Z^Ss and ZnaC^. Another way to check up the correctness of a formula is to repre- sent each combining power by a line drawn from the symbol of the element; the number of lines will, evidently, be equal to the valence of the element. A number of formulas written in this way will be self-explanatory: ^ Na x AlC >0 Zn = Na/ Formulas written in the way illustrated above are called graphic formulas; they are useful for other purposes than the one explained here. In representing compounds by graphic formulas it is impor- tant to emphasize the fact that there can be no free combining powers the lines must run from the symbol of one element to that of another. Such combinations as the following do not rep- resent compounds: Na, / ">O AJ = 120 INORGANIC CHEMISTRY FOR COLLEGES From the above it will be seen that much mental effort will be saved if the valence of an element is learned as soon as its com- pounds are met with, for if this is done it will not be necessary to memorize a large number of formulas. 134. The valence of certain elements is constant; thus, hydrogen, sodium, and potassium always have the valence 1; zinc, calcium, and magnesium, 2; aluminium, 3, etc. The valence of other elements varies ; iron, for example, has the valence 2 in some compounds and 3 in others ; it forms two chlorides which have the formulas FeCk and FeCla, respectively. This is a source of complexity, but one that is readily mastered as the compounds are studied. 135. The method of writing formulas of compounds as de- scribed above is applicable in the case of those which contain two elements only. It can be applied in a simple way to other com- pounds when a few principles are made clear. We have seen that the formula of sulphuric acid is B^SCU and that of sodium sulphate is Na2S(>4. There are many substances known which can be pre- pared by replacing the hydrogen in sulphuric acid by atoms of metals; all these compounds contain the group of atoms repre- sented by the symbols SO 4; they are all called sulphates. We have calcium sulphate CaSCU, copper sulphate, CuSO4, zinc sul- phate, ZnSC>4, etc. The group of atoms present in these com- pounds, SO4, is called a radical, the word being derived from the Latin word radix, meaning root. As it is present in sulphuric acid it is called an acid radical. We can assign to this radical a valence; since it is found in combination with 2 hydrogen atoms in sul- phuric acid, H2SO4, its valence is 2. With this understood it is possible to write the formula of the sulphate of any metal pro- vided the valence of the latter is known. The method described in detail above when applied to writing the formulas of sodium sulphate, zinc sulphate, and aluminium sulphate leads to the fol- lowing: i ii ii ii in ii Na 2 S0 4 ZnSO 4 A1 2 (SO 4 )3 In this case the valence of the SO4 radical is written above it, in the way used with the elements. In the formula for aluminium sulphate the radical is enclosed in a parenthesis and a subscript 3 is written to the right of it; this indicates that 3 radicals each with CHLORINE. VALENCE 121 the valence 2, 3 X 2 = 6, are needed to combine with 2 atoms of aluminium with the valence 3, 2 X 3 = 6. In order to write the formulas of the compounds derived from acids by replacing the hydrogen of the latter by metals compounds which are called salts it is necessary to know the valences of the acid radicals. EXERCISES 1. Chlorine can be prepared by the action of hydrochloric acid on lead dioxide, PbO 2 , and on potassium permanganate, KMnO 4 . Convert the following into balanced chemical equations: (a) PbO 2 + HC1 > PbCl 2 + H 2 O + C1 2 and (b) KMnO 4 + HC1 -+ KC1 + MnCl 2 + H 2 O + C1 2 . 2. (a) What proportion of the chlorine in sodium chloride is set free when the latter is treated with sulphuric acid and manganese dioxide? (b) What proportion of the chlorine is set free when hydrochloric acid and manganese dioxide are used? (c) What determines which is the cheaper process? 3. The labels on cans of bleaching powder usually contain a statement as to the available chlorine present in the powder. Available chlorine is that which is set free when bleaching powder is treated with an acid. If the amount of this chlorine is stated as 35 per cent, what is the percentage purity of the powder? 4. How could you distinguish by chemical means the folio wing: (a) CaOCl 2 and CaCl 2 , (b) NaCl and Na 2 SO 4 , (c) MnO 2 and C? 5. Hydrogen is a very light gas and chlorine is very heavy. If a cylinder containing the former is inverted and placed above one containing chlorine, the two gases will slowly mix and, after a time, they will be uniformly dis- tributed. Can you suggest any hypothesis to explain this fact? 6. If 100 liters of liquid chlorine were allowed to change to a gas what volume would it occupy at and 760 mm.? (b) If the gas were mixed with air so that it was present in the proportion of 1 part to 1,000,000 of air what would be the volume of the mixture? (c) Assuming that the mixture formed a layer 10 meters high what area would it cover? 7. (a) Into what compound did the large quantities of chlorine liberated in the recent war probably change? (6) What becomes of the chlorine that escapes in a chemical laboratory? (c) Why is chlorine-water kept in bottles of brown glass? 8. Can you think of any use in warfare that might be made of the fact that chlorine reacts with certain hydrocarbons to form carbon and hydro- chloric acid? The reaction in the case of acetylene is practically instantaneous. 9. Chlorine is both a disinfectant and a deodorizer. Why? 10. How could you find out if a sample of hydrogen chloride contained a small amount of chlorine, if there were not enough of the latter present to be distinguished by its color? 11. Which of the substances indicated by the following formulas would 122 INORGANIC CHEMISTRY FOR COLLEGES you expect to be inflammable? (a) CO 2 , (6) CC1 4 , (c) CHC1 3 , (d) CH 4 , (e) PH 3 , (/) P 2 5 , (<7) C 6 H 6 , (h) C 2 H 6 O, (i) NO 2 . 12. Convert the following into balanced equations: (a) C^He + O * CO 2 + H 2 0; (6) SbCl 3 + H 2 O -> Sb 2 O 3 + HC1; (c) KC1 + H 2 SO 4 - HC1 + + K 2 SO 4 ; (d) CH 4 + Br 2 -> CHBr 3 + HBr; (c) C 2 H 6 O + O 2 -> CO 2 + H 2 O. 13. The element bromine forms a compound of the formula HBr: (a) What is the valence of the element? (6) Write the formulas of the com- i ii in in n pounds of bromine with Na, Ca, Al, Sb, and Zn. (c) What is the valence of sulphur in the compound ZnS? (d) Write the formulas of the compounds of sulphur and Na, Ca, and Sb. (e) What is the valence of the SO 3 radical in H 2 SO 3 ? (/) Write the formulas of compounds containing this radical i n and K and Ca. (g) The formula of nitric acid is HNO 3 , write the formulas i n in of the compounds containing the NO 3 radical and K, Cu, and Al. 14. At 100 volumes of water dissolve 461 volumes of chlorine. Cal- culate the percentage of chlorine by weight in the mixture. 15. Calculate the weight of 1 liter of chlorine, C1 2 , at and 760 mm. from its atomic weight. 16. If 2 liters of chlorine are dissolved in 1 liter of water and the mixture exposed to the sunlight so that the oxygen formed is collected, what volume of the latter would be obtained? CHAPTER XI HYDROCHLORIC ACID. DOUBLE DECOMPOSITION 136. The study of hydrochloric acid will furnish an opportunity to learn a great deal about a very important class of substances, called acids, which possess many properties in common. We come in contact often in daily life with compounds of this class. Acids are formed when many organic substances undergo decay through the influence of the bacteria present in the air; the juices of fruits and vegetables ferment, and gases and acids are formed; and milk turns sour. Hydrochloric acid is of particular interest because it plays a very important part in the processes that take place in the animal body. The digestion of certain foods in the stomach is brought about through the action of hydrochloric acid and a substance called pepsin, both of which are present in gastric juice. Pepsin, which can be obtained for experimental purposes from the lining of a pig's stomach, does not, alone, bring about the digestion of foods, but in the presence of hydrochloric acid becomes very active; the latter acts probably as a catalytic agent. As the hydrochloric acid enters into combination with other substances and is excreted from the body, some compound which will serve as a source of this acid must be taken with the food. Chlorides are present in many foods; as these, however, do not generally furnish an adequate supply of the acid, common salt, sodium chloride, is an important constituent of the diet. Animals have been known to risk death in a desire to get salt. Hydrochloric acid is frequently prescribed by physicians in those cases of dys- pepsia which are supposed to be due to the lack of the acid in the gastric juice. 137. Historical. Hydrochloric acid was known to the early alchemists, who probably prepared it by heating a mixture of salt and green vitriol (FeSO4,7H 2 O). Glauber, one of the most 123 124 INORGANIC CHEMISTRY FOR COLLEGES famous of the alchemists, described in 1648 the preparation of the acid from salt and sulphuric acid. The other product formed in the reaction, sodium sulphate, Na2S04,10H2O, was for a long time known as Glauber's salt. Priestley was the first to obtain the acid free from water, and collected the gas over mercury. 138. Preparation of Hydrochloric Acid. The acid can be made by burning hydrogen in chlorine a reaction which has been discussed in some detail (120). The equation for the reaction is repeated here: H 2 + C1 2 = 2HC1 This is not ordinarily a practical method because it involves the previous preparation of both hydrogen and chlorine. It is used, however, to manufacture hydrochloric acid when a very pure acid is required, and when hydrogen and chlorine are available. This is the case when sodium hydroxide is made by the electrolysis of a solution of sodium chloride. The method used in the laboratory and on the large scale is based on the reaction between salt and concentrated sulphuric acid. When the two substances are heated together at a high temperature a reaction indicated by the following equation takes place : 2NaCl + H 2 SO 4 = Na 2 SO 4 + 2HC1 At lower temperatures those used in the laboratory the sub- stances interact in different proportions: NaCl + H 2 SO 4 = NaHS0 4 + HC1 The products of the reaction in the first case are sodium sulphate and hydrochloric acid; in the second case sodium hydrogen sulphate is formed. The acid escapes as a gas and can be col- lected by the upward displacement of air, since it is heavier than air, or by conducting it into water in which it dissolves. 139. Physical Properties of Hydrochloric Acid. The acid is a colorless gas, which is heavier than air; 1 liter at and 760 mm. pressure weighs 1.641 grams. One liter of water at dissolves 525 liters of the gas. In order to distinguish between the pure compound and its solution in water, the former is sometimes called hydrogen chloride. HYDROCHLORIC ACID. DOUBLE DECOMPOSITION 125 The marked tendency of hydrogen chloride to dissolve in water can be shown by a striking experiment. A large glass globe is filled with the dry gas and then closed by a stopper carrying a long glass tube and a medicine dropper filled with water and so placed that by means of it a small amount of water can be forced into the globe (see Fig. 15). The apparatus is then inverted and supported over a reservoir con- taining water colored by blue litmus. When water is forced into the globe by pressing the rubber end of the medicine dropper, the gas dissolves in it; this leaves a partial vacuum, and the water in the reservoir rushes up to take the place of the dissolved gas; as it enters, the blue litmus is changed to red by the acid. 140. Chemical Conduct of Hydro- chloric Acid. When the gas is dry it is very inactive, but a small amount of water serves as a catalytic agent in helping to bring about reaction be- tween it and other substances. Dry hydrochloric acid does not react with zinc, even when the gas is dissolved in a liquid such as benzene. The addition of a small amount of water, however, causes a reaction to begin, and hydrogen is evolved. We have learned that acids change the color of litmus paper from blue to red. If a dry piece of blue litmus paper is put into the dry gas, it is scarcely affected; if the paper is first moistened it turns red immediately. Hydrochloric acid as ordinarily prepared contains enough water to catalyze the reactions between it and other substances. Under these circumstances it reacts with ammonia. This can be shown by a striking experiment. A large cylinder, the open end of which is carefully ground flat and covered with vaseline, is filled with the gas by upward displacement of air. It is then closed by placing over it a ground- glass plate, the vaseline serving to make the covering air-tight. A similar cylinder of ammonia, which is also a colorless gas having the formula NHs, is prepared and covered. FIG. 15. 126 INORGANIC CHEMISTRY FOR COLLEGES The cylinders are now placed mouth to mouth, the one containing ammonia on top. The two glass plates are next quickly with- drawn and the open ends of the cylinders brought into close con- tact. A dense white cloud appears at once, formed as the result of the combination of the two invisible gases; the substance is ammonium chloride, which is often called sal ammoniac, and is used in certain kinds of electric batteries. The reaction is ex- pressed by the following equation: HC1 + NH 3 = NH 4 C1 If an attempt is made to pull apart the cylinders, it will be found that considerable force must be exerted. This is because the gases have disappeared, and in separating the cylinders we must overcome the pressure of the atmosphere on them, which is not appreciated as long as they were filled with gas. The cylinders can be separated easily by sliding one over the other, since in doing this we are not working against the air-pressure. When a small opening is made in this way the air rushes in to fill the vacuum. 141. When a jar of hydrochloric acid is opened to the air a white cloud is observed at the mouth of the jar. The cloud is denser if the breath is blown through the gas. Water can exist in the air in the form of an invisible gas, called water-vapor. When certain substances which are very soluble in water are brought in contact with moist air, they absorb water from the latter, and, finally, a solution of the compound is formed. The cloud produced from hydrochloric acid consists of minute drops of a solution of the gas in water. When a gas behaves in this way, it is said to fume in the air. Solids which take up moisture from air, become damp, and finally pass into solution in the water absorbed, are said to be deliquescent. Calcium chloride, Cadi, is an example of such a salt. 142. Properties of Aqueous Solutions of Hydrochloric Acid. When hydrogen chloride is dissolved in water, the solution exhibits many important properties which are not shown by the anhy- drous compound. The word anhydrous means free from water; it is a very convenient adjective and is often used in chemistry. The solution of hydrochloric acid in water, like the solutions of other acids, has a sour taste. It turns blue litmus paper red, and effects a change in color in other substances. The dye called HYDROCHLORIC ACID. DOUBLE DECOMPOSITION 127 methyl orange is yellow ; it is changed to a red compound when treated with an acid. The salts of phenolphthalein are red, but are converted into a colorless compound by acids. Substances which behave in this way are called indicators; they are much used in chemistry to detect the presence of acids. Hydrochloric acid reacts with the more active metals and, as a result, hydrogen and a chloride are formed. A number of such reactions have already been discussed (46). Two typical cases are represented by the following equations: Zn + 2HC1 = ZnCl 2 + H 2 2A1 + 6HC1 = 2A1C1 3 + 3H 2 Hydrochloric acid reacts with the oxides of metals: Na 2 + 2HC1 = 2NaCl + H 2 FeO + 2HC1 = FeCl 2 + H 2 O A1 2 3 + 6HC1 = 2A1C1 3 -h 3H 2 The acid reacts also with hydroxides : NaOH + HC1 = NaCl + H 2 Zn(OH) 2 + 2HC1 = ZnCl 2 + 2H 2 O A1(OH) 3 + 3HC1 = A1C1 3 + 3H 2 All acids enter into reactions similar to those illustrated by means of hydrochloric acid in the three series of equations just given. 143. Commercial hydrochloric acid, which is often called muri- atic acid, possesses a yellow color which is caused by the presence of impurities in it. It usually contains small quantities of sulphuric acid, chlorine, iron chloride, and arsenic. It contains about 40 per cent by weight of hydrogen chloride, the rest being chiefly water. It is sold in large glass vessels called carboys; each vessel is enclosed in a wooden box and is packed in straw. Carboys for shipment of acid generally hold about 12 gallons. A purer quality of the acid can be purchased: this is sold under the designation C.P., the letters being an abbreviation of the words chemically pure. So-called C.P. chemicals are supposed to contain not more than traces of impurities, the amounts of these present not being enough 128 INORGANIC CHEMISTRY FOR COLLEGES to interfere with the use of the chemical for all ordinary purposes. It takes a very long time to obtain any substance in such a pure condition that the presence of other substances in minute amounts cannot be detected in it by a skilled chemist. Hydrochloric acid, C.P., contains 40 per cent by weight of HC1 and has the specific gravity 1.2, that is, it is 1.2 times as heavy as an equal volume of water. Acid of this strength is called concentrated hydrochloric acid; it is made by passing the gas into water as long as it is absorbed at ordinary temperatures; such a solution is said to be saturated. Dilute hydrochloric acid can be made of any strength less than this; that commonly supplied in the laboratory contains about 20 per cent of HC1. What happens when hydrochloric acid is heated is determined by the proportion of hydrogen chloride it contains. If the solution is concentrated, the gas is driven off and if dilute, water-vapor first escapes. In either case as the acid is heated a point is finally reached when the product that distills over boils at 110 at 760 mm., and contains 20.24 per cent HC1, and 79.76 per cent EkO. This is called the constant-boiling mix- ture of water and hydrochloric acid. Hydrochloric acid is used in the preparation of chlorine and of chlorides, in cleaning metals, and for various purposes in the chemical laboratory. 144. Chlorides. Chlorides can be formed by the action of chlorine on the various elements; a number of examples have already been noted (121). In the case of the metals, chlorides can be prepared by means of the reactions which were illustrated by the equations given in the last section. Many metals which do not react with hydrochloric acid form oxides which are converted into chlorides by the acid. For example, copper and hydrochloric acid do not react, but copper oxide dissolves readily in the acid: CuO + 2HC1 = CuCl 2 + H 2 O The uses to which chlorides are put will be described in connection with the consideration of the metals which they contain. 145. Test for Chlorides. It is often necessary to determine whether any particular substance is a chloride, or whether a mix- ture contains a chloride. Such problems arise in what is known as qualitative analysis a branch of chemistry which has to do with HYDROCHLORIC ACID. DOUBLE DECOMPOSITION 129 the determination of the presence of the various substances in com- mercial and other products. We turn to the chemist to analyze a supposed poison, a sample of steel or brass, or a patent medicine; to test milk and drinking water for impurities, or a rock for gold. All such problems are solved by qualitative analysis. The pres- ence of any substance in a mixture is established by separating and identifying it, or by converting it into another substance the properties of which are known. The method used can be illus- trated clearly by a consideration of the way in which the presence of a chloride is determined. In this case it is not necessary to separate the chloride itself, provided all we want to know is whether a chloride is present. The substance to be tested is dissolved in water; if it is insoluble in water another liquid is used. A solu- tion of silver nitrate is next added. If a chloride is present a white cloud appears or a white solid separates and settles to the bottom of the tube. When an insoluble substance separates in this way on mixing two solutions, it is called a precipitate. The substance in this case is silver chloride; the reaction is represented by the following equation when the chloride present is sodium chloride: NaCl + AgNO 3 = NaNO 3 + AgCl It is seen that the metals change places, and we get as the result of the reaction the chloride of silver and the nitrate of the metal which was originally in combination with chlorine. The chlorides of the metals in general act in this way, for example: ZnCl 2 + 2AgNO 3 = Zn(NO 3 ) 2 + 2AgCl A1C1 3 + 3AgN0 3 = A1(NO 3 ) 3 + 3AgCl We could use the formation of the white precipitate when silver nitrate is added to the solution, as a test for chloride, provided no other substances produced a similar result. Other substances, however, do give a white precipitate under these conditions, and we must find some property of silver chloride other than its color and insolubility in water, if we are to use its formation to show the presence of chlorides. Silver chloride is insoluble in dilute nitric acid, whereas the white precipitates formed when substances other than chlorides are treated with silver nitrate are soluble in this acid. This important fact makes it possible to use the reaction 130 INORGANIC CHEMISTRY FOR COLLEGES with silver nitrate as a test for a chloride; we add dilute nitric acid to the solution before adding silver nitrate, and if a white precipitate is formed we conclude that a chloride is present. This method of detecting a chloride in the presence of other sub- stances is applicable in the case of practically all commonly occur- ring compounds ; it is used as a test for chlorides. 146. When we wish to test for a substance or one of its con- stituents, we convert it into a new substance possessing charac- teristic properties which can be readily recognized, and the sum of which are not possessed by any other substance that could be formed under the conditions used in the test. We use at times the production of an odor, a color, or an insoluble substance, the forma- tion of the latter being accomplished if possible. In learning a test the student should take care to understand what are the characteristic properties of the substances upon the formation of which the test is based. The test for a chloride is based on the fact that silver chloride is a white solid which is insoluble in nitric acid. Since silver bromide is insoluble in nitric acid and possesses a faint yellow color, it might be mistaken by the beginner for silver chloride; to avoid this, a solution of ammonia is added to the pre- cipitate; silver chloride dissolves readily, whereas silver bromide requires a large amount of the solution to dissolve it. It should be noted here, also, that silver chloride when exposed to sunlight soon turns in color to a light violet shade; silver bromide does not do this. In order to be sure that the conclusion drawn from a test is correct, it is advisable for the beginner to apply independent tests, if possible, to confirm his conclusion. A test made for this purpose is called a confirmatory test. Accumulation of evidence in support of a statement makes it more worthy of belief; we cannot be too careful in drawing our conclusions. 147. Types of Chemical Reactions. A large number of chem- ical reactions have been discussed up to this point. It is advisable to reexamine some of them and see how they can be put into a few simple classes. One of the chief aims of science is to systematize knowledge, and as has been pointed out, the student should con- stantly attempt to correlate the new facts as they are presented, that is, to see their relation to facts already mastered. The simplest type of reaction which has been studied is that which occurs when two substances unite to form a third substance; HYDROCHLORIC ACID. DOUBLE DECOMPOSITION 131 this is called combination and is illustrated by many equations already given, for example: 2Cu + O 2 = 2CuO The reverse of combination is decomposition: 2HgO = 2Hg + 2 Displacement is a third type of reaction: Zn + H 2 SO 4 = ZnS0 4 + H 2 In the above case hydrogen has been substituted for zinc in H 2 S0 4 . 148. A fourth type is illustrated by such a reaction as the following: NaCl + AgN0 3 = AgCl + NaN0 3 In this case since two substances interact and form two new sub- stances, the reaction is called one of double decomposition or meta- thesis. A reaction is put into this class only when the two sub- stances interact in a particular way, which will be explained in detail, since reactions of double decomposition are examples, per- haps, of the most important type that the student will meet. The significance of the expression acid radical has been explained; it is the group of atoms other than hydrogen present in the acids. For example, the radicals present in the acids having the formulas HC1, HNO 3 , H 2 SO 4 , H 2 SO 3 , H 3 PO 4 , are Cl, N0 3 , SO 4 , SO 3 , and PO 4 , respectively. Compounds made up of metallic atoms and acid radicals are called salts, for example, AgNO 3 , Na 2 SO 4 , etc. It will also be recalled that bases are com- pounds which contain metallic atoms in combination with one or more hydroxyl groups, such as NaOH, Zn(OH) 2 , etc.; the hydroxyl group has the valence 1. Reactions of double decomposition take place between pairs of compounds belonging to these classes. With these facts before us we can examine fully the reactions of this type which have been studied, and become prepared to under- stand those that are to follow. An examination of the equation last given, NaCl + AgN0 3 = AgCl + NaN0 3 , brings out the fact that the reaction is between two salts, and that the two products of the reaction are formed as the result of an 132 INORGANIC CHEMISTRY FOR COLLEGES exchange of the metallic atoms; sodium takes the place of silver, and silver that of sodium. A much better name than double decomposition for this class of reaction would be double sub- stitution. We have seen that metallic atoms can take the place of hydrogen in acids; an example of this kind of reaction is given above as an example of substitution. Such an exchange is possible in double decomposition reactions; the formation of hydrogen chloride from salt and sulphuric acid is an example: 2NaCl + H 2 SO 4 = Na 2 S0 4 + 2HC1 In this case a salt and an acid enter into double decomposition. It is seen from the two examples given that a double decomposition takes place between two substances when their hydrogen or metallic atoms mutually replace each other. A double decomposition occurs also when only single atoms exchange places, for example, NaCl + H 2 S0 4 = NaHS0 4 + HC1 In this case but 1 hydrogen of the sulphuric acid changed place with 1 sodium atom. Cases of this kind are not so common as those in which all the atoms united with the acid radical enter into the exchange. A few double decompositions will be considered in detail to illustrate how reactions of this type can be written if we apply the statement just given. Suppose it is desired to represent a double decomposition between the compounds sodium chloride and silver sulphate. We first set down the formulas of these compounds and in order to indicate how the molecules break apart draw dotted lines between the metallic or hydrogen atoms and the acid radicals; we also indicate their valencies thus: ill i i II Na | Cl + A g2 | S0 4 We then change the positions of the metallic atoms, neglecting the subscript numbers representing the number of atoms involved in everything except the acid radicals, thus : I i I I ! II ill I ! II Na ! Cl + Ag 2 I S0 4 - Ag | C1 + Na ! S0 4 i I i i HYDROCHLORIC ACID. DOUBLE DECOMPOSITION 133 We next add subscript numbers which are necessary to satisfy the valence requirement. In the above case a 2 must be placed to the right and below the symbol for sodium in sodium sulphate in order that the salt may have the correct formula. Finally the equation is balanced ; this is done by taking two molecules of NaCl and two of AgCl. The equation becomes as a result: 2NaCl + Ag 2 SO 4 = 2AgCl + Na 2 SO 4 Another double decomposition will be written in this way between aluminium chloride and sulphuric acid. First represent the parts which will interchange and the result of the interchange, affixing the numbers representing the valence of each, in i i i j ii i i i in ; ii Al ! C1 3 + H 2 ! SO 4 -* H ! Cl + Al I S0 4 i i ! ! Next add the subscript numbers to satisfy the valence require- ments : in i i i i ii iii in ! ii Al | C1 3 + H 2 I SO 4 -> H ! Cl + A1 2 ! (SO 4 ) 3 | j I I Finally balance the equation: 2A1C1 3 + 3H 2 S0 4 = 6HC1 + A1 2 (S0 4 ) 3 All these operations can be done in successive steps without writing the equation three times as was done here to make each step clear. Double decompositions which involve bases are written in the same way: I i I III I I III ! I I i I Na | OH + Al j Cla - Al | OH + Na I Cl ill i 3NaOH + A1C1 3 = A1(OH) 3 + 3NaCl In double decompositions between acids and bases water is formed : Na j OH + H 2 | S0 4 = Na | SO 4 + H | OH 2NaOH + H 2 SO 4 = Na 2 S0 4 + 2H 2 O From the above considerations we see that double decompositions can be written between the following pairs of substances: Acids 134 INORGANIC CHEMISTRY FOR COLLEGES and salts, acids and bases, bases and salts, and salts and salts. It is right to ask at this point, do all the reactions written in this more or less mechanical way represent chemical changes that actually take place? It has been already emphasized that it is impossible to arrive at a chemical formula or reaction by a simple mathe- matical juggling of symbols. We shall see later that all the chem- ical changes indicated by double decomposition reactions between the pairs of substances listed, do take place to a greater or less extent. Only when the reaction is practically complete in the manner indicated by the equations do we call them reactions of double decomposition. 149. A study of many reactions of this type has brought out the important generalization that if one of the products formed can escape in any way, the double decomposition takes place. When, for example, a gas is formed, it leaves the vessel in which the reac- tion is taking place. The reaction also takes place if one of the products is insoluble; it precipitates and is no longer in solution where it can react; it leaves the field of action. We can readily see why this removal of one of the products affects the course of a reaction. Double decompositions belong to an important class of reactions known as reversible reactions. In reactions of this kind the substances indicated on the left-hand side of the equation can interact to form those indicated on the right-hand side, and those on the right can form those on the left. We express this in the equation for the reaction by replacing the sign of equality by another symbol, thus : NaCl + H 2 SO 4 <= NaHSO 4 + HC1 Written in this way, the equation indicates that sodium chloride and sulphuric acid can react to form sodium hydrogen sulphate and hydrochloric acid, and, also, that sodium hydrogen sulphate and hydrochloric acid can react to form sodium chloride and sul- phuric acid. It might well be asked, under what circumstances does the reaction take place in one direction only. It is evident that if the hydrochloric acid is taken away as soon as it is formed, it cannot react with the sodium hydrogen sulphate produced with it to form sodium chloride again. We are now in a position to understand under what circumstances reversible reactions of this kind proceed to completion in one direction, and become, thus, HYDROCHLORIC ACID. DOUBLE DECOMPOSITION 135 reactions of double decomposition. This occurs when one of the products is removed, the removal taking place of itself if a gas which escapes is formed, or if an insoluble substance is produced which separates as a precipitate. With this knowledge we can examine any reaction written between the pairs of substances listed above, and state whether they represent double decomposition reactions. In order to do this it is necessary to know whether either of the substances produced is a gas which escapes or an insoluble compound. Double decomposition also takes place when water is one of the products of the reaction; the reason for this will be explained fully later. In brief, then, reactions of double decomposition take place between acids, bases, and salts when one of the products is a gas, an insoluble substance, or water. This is a generalization of the greatest importance, because a large proportion of the reactions to be met with belong to this type, EXERCISES 1. Calculate from the formula of hydrochloric acid the weight of 1 liter of the gas at and 760 mm. 2. (a) What is meant by the expression "water seeks its own level"? (6) What causes the water to rise into the globe in the experiment described in section 139? (c) How would you tell after the experiment how much air was mixed with the hydrogen chloride in the globe? 3. (a) Why does concentrated hydrochloric acid fume in the air? (6) Why does it fume more strongly if the breath is blown across the open mouth of the bottle? 4. Write equations for the reactions which take place between hydro- chloric acid and the following: (a) ZnO, (6) K 2 O, (c) Fe, (d) Fe 2 O 3 , (e) Sn, (/) Mg, (g) Fe(OH) 2 , (h) A1(OH) 3 . 5. (a) Write equations for three different reactions by which ZnClc can be made. (6) Can CuCl2 be made by similar reactions? 6. Write equations for reactions involved when the following are tested to determine whether they are chlorides : HC1, CaCl 2 , FeCl 3 . 7. (a) Is silver nitrate necessary if it is desired to test for a chloride? (6) If not, what other salt could be used? (c) Write equations for the reac- tions which would take place if this salt were used in testing solutions of NaCl, ZnCl 2 , and A1C1 3 . 8. (a) How much hydrochloric acid, HC1, is required to dissolve 10 grams of zinc? (b) If concentrated hydrochloric acid which contains 40 per cent HC1 is used, how much of it is necessary to dissolve the zinc? (c) Acid of this strength has the specific gravity 1.2. What volume of the acid is required? 136 INORGANIC CHEMISTRY FOR COLLEGES 9. How many grams of (a) concentrated hydrochloric acid and (b) the constant-boiling mixture of HC1 and water must be taken to obtain 1 gram- molecular-weight of HC1? (c) With what weight of sodium hydroxide will this amount of acid react? (d) If this weight of acid is mixed with enough water to make the volume of the mixture 1 liter, how much sodium hydroxide is present in a solution which reacts with 50 c.c. of the solution of the acid? 10. Using the method described in the text, write reactions of double decom- position between the following pairs of compounds: (a) NaOH and HC1, (6) KOH and H 2 SO 4 , (c) NaOH and H 3 PO 4 , (d) CaCl 2 and Na 2 SO 4 , (e) A1C1 3 and KOH, (/) A1C1 3 and H 2 SO 4 , (g) CaCl 2 and Na 2 PO 4 . 11. What weight of silver chloride, AgCl, is formed when 10 grams of silver nitrate, AgNOs, are treated with hydrochloric acid? 12. A piece of a silver dime weighing 0.200 gram was dissolved in nitric acid and the solution of silver nitrate, AgNO>, formed was treated with hydro- chloric acid. The silver chloride which was precipitated was dried and found to weigh 0.239 gram, (a) How much silver was present in this weight of silver chloride? (b) What percentage of silver did the dime contain? 13. What weight of sodium chloride is required to furnish enough hydrogen chloride to fill at and 760 mm. a flask of 5 liters capacity? CHAPTER XII THE ENERGY FACTOR IN CHEMICAL CHANGE 150. The important fact has already been emphasized that in all chemical changes the transformations brought about in mat- ter are associated with changes in chemical energy. Energy manifests itself through the effect it produces on matter; these effects alone appeal to our senses, but we must study their cause if we are to get at a fundamental conception of matter itself. Chemical energy cannot be directly measured; we determine the amount of this form of energy involved in a change by transforming it into other kinds of energy which can be measured. It has already been pointed out that one form of energy can be changed into another form (12). When this fact was studied quanti- tatively it was discovered that in the transformations no energy was lost or gained. A number of investigators announced inde- pendently what is known as the law of the conservation of energy (J. R. Mayer in 1842 and Helmholtz in 1847). The law states that within a system undergoing change there is no loss or gain in energy. In certain cases one kind of energy can be com- pletely transformed into another kind; whereas in other cases the transformation brings about the appearance of two kinds of energy. Electrical energy can be completely converted into heat, but when an attempt is made to change it into light, a large part of the energy is transformed into heat. Changes of the first kind are used in measuring energy. 151. When two substances enter into a chemical reaction, chemical energy is changed into other forms of energy. It is possible to carry out these reactions in such a way that heat is the only form of energy produced. Since the amount of heat can be measured, the total change in energy that accompanies a chemical reaction can be determined. We make use of such 137 138 INORGANIC CHEMISTRY FOR COLLEGES measurements, for example, in comparing the values of coal, gas, alcohol, etc., as sources of heat when they are burned. In comparing the chemical behavior of two substances, for example, the behavior of iron and of silver toward chlorine, we are concerned with the relative tendencies of the two metals to enter into the reaction involved. The tendency of one element to react with another to form a compound is not measured by the heat produced when the reaction takes place, but by the work necessary to reverse the chemical change. For example, in the case of the formation of silver chloride from silver and chlorine the tendency for the compound to form is measured by the work required to separate one gram-molecular-weight of the chloride into silver and chlorine at the pressure of one atmosphere. It is evident that the greater the tendency for a compound to form, the greater must be the work required to separate it into its constituents. The work involved in such separations can be measured in several ways. One of these involves the determination of the amount of electrical energy required to effect the decomposition of the compound into the elements from which it was formed. For example, we can measure the amount of electrical energy required to separate one gram-molecular-weight of silver chloride, into silver and chlorine at the pressure of one atmosphere. This work is taken as the measure of the tendency of silver and chlorine to unite. 152. When chemical changes take place, heat is produced or disappears. The importance of this fact was recognized early, and measurements of the heat change in many chemical reactions have been made. Although these heat values are not a measure of the tendencies of the reactions to take place, they are helpful, however, in comparing the relative tendencies in the case of similar reactions. It has been found, for example, that the tendencies of the several metals to unit with chlorine are in the same order as the values of the heat produced when the chlorides of these metals are formed. 153. Measurement of Heat Energy. We must next turn our attention to the way in which heat energy is measured ; electrical energy will be considered later. We are familiar with the thermometer and know that it is used to measure temperature; it tells us how hot a thing is but it gives THE ENERGY FACTOR IN CHEMICAL CHANGE 139 us no indication of how much heat is present. If we were to test the flame of a burning match we should find that it was as hot as the flame from a log of wood ; the two would show approximately the same temperature, but with one we could heat a room and with the other we could not ; the quantity of heat is different in the two cases. We see, thus, that heat energy has two factors intensity and quantity. We measure intensity by means of a thermometer, and quantity with what is called a calorimeter. The amount of heat is measured by the rise in temperature which the heat pro- duces in a definite weight of water. The unit in this case is called a calorie; it is the amount of heat required to raise the temperature of one gram of water one degree cen- tigrade (usually from 15 to 16). In technical work a unit based on the English system of weights is used; it is called a British thermal unit (B.t.u.) and is the amount of heat required to raise one pound of water one degree Fahrenheit; 1 B.t.u. =252 calories. If it is desired to know how much heat is generated when a certain re- action takes place, weighed quantities of the substances involved are placed in an apparatus surrounded by water, the temperature of which is known. The substances are next allowed to react, and as a result the heat generated causes the temperature of the water to rise. If this rise is noted and the weight of the water is known, the number of calories liberated by the reaction can be calculated. The apparatus in which such observations are made is called a calorimeter. A convenient form that is often used is represented by Fig. 16; it is called a bomb-calorimeter because chemical reactions under pressure can be carried out in it. Sup- pose it is desired to know how much heat is generated when char- coal burns. A weighed amount of charcoal is put into the cup, the bomb is then closed and oxygen is forced in under pressure. The reaction is started by passing a current of electricity through FIG. 16. 140 INORGANIC CHEMISTRY FOR COLLEGES the wire which passes over the charcoal. The wire melts and ignites the charcoal. A calorimeter of this type is used to determine the heat pro- duced when foodstuffs are burned. The value of a food is deter- mined, in part, by the amount of heat which it produces when oxi- dized, since one function of food is to furnish heat to keep up the temperature of the body. One gram of fat gives about 9000 cal- ories when it burns, whereas 1 gram of starch gives about 4000 calories. Such facts as these are of vital importance in the science of foods. 154. Thermochemistry. In all substances as we know them, matter is associated with energy. Up to this point the changes in the matter have been emphasized, but for a more complete under- standing of these changes we must study the transformations in energy which take place simultaneously with the changes in matter. Two substances possessing distinctly different properties maybe composed of the same kind of matter. A case which will be dis- . cussed in detail later can be mentioned here. A diamond is a very different substance from a bit of charcoal, but the matter in the two is identical they consist solely of the element called carbon. The great difference in their properties can be traced to the fact that the amount of energy which is combined with the elementary substance carbon is different in the two cases. This can be shown by a careful study of what occurs when the two substances are burned. If 12 grams of charcoal are burned, 44 grams of the gas carbon dioxide are obtained, the equation for the reaction being C + (>2 = CO2. If 12 grams of diamond are burned, 44 grams of carbon dioxide are also obtained; the gases formed in the two cases are identical in weight and properties. The study of the matter involved in the two cases gives us no information as to the cause of the difference between charcoal and diamond. If the heat produced in the two cases is determined if the energy change is investigated we find that the results are different. When 1 gram of charcoal burns the heat produced is 8080 calories; when 1 gram of diamond is burned 7860 calories are set free. Since the same amount of oxygen is transformed into carbon dioxide in the two cases, we must attribute the different results in the two cases to the fact that carbon as charcoal contains gram for gram more chemical energy than diamond. The element carbon is in THE ENERGY FACTOR IN CHEMICAL CHANGE 141 both these substances. We must differentiate between the element carbon, and the elementary substance carbon, which can appear in a variety of physical forms as the result of the fact that the element can be associated with different amounts of chemical energy. Carbon is present in carbon dioxide, in wood, in sugar, in alcohol, and in thousands of other compounds; it has no characteristic property except weight. 155. An addition can be made to chemical equations to express the amount of heat set free when the reaction indicated by the equation takes place. For example, the burning of charcoal in oxygen can be represented thus: C + O 2 = CO 2 + 97,000 cal. Such an expression is a thermochemical equation. In this case C signifies the number of grams of carbon equal to its atomic weight, that is, 12 grams. Each symbol in general represents the weight in grams equal to its atomic weight; this weight is called a gram- atomic-weight or gram-atom. In a thermochemical equation the formula CC>2 stands for 12 + (2 X 16) = 44 grams of carbon dioxide; this is a gram-molecular-weight or a mol of carbon diox- ide. The above equation means, then, that when 12 grams of carbon burn in oxygen, 32 grams of the latter unite with it and form 44 grams of carbon dioxide, and, at the same time, 97,000 calories are set free; this number of calories is called the heat of formation of carbon dioxide. A chemical reaction in which heat is given off is said to be exothermic; one in which heat is absorbed is endothermic. When an exothermic compound breaks down into its elements the energy absorbed is equivalent in amount to that given off in its formation; in this case heat is converted into chemical energy. Thermochemical equations deal with but one factor in the energy changes which take place in chemical reactions the quantity factor. For a complete understanding of these changes a knowledge of the intensity factor is necessary. We shall see later that the transformation of chemical energy into electrical energy will give us some knowledge of this intensity factor. A few of the equations which have already been given will be discussed briefly from the standpoint of energy. The thermo- 142 INORGANIC CHEMISTRY FOR COLLEGES chemical equation for the burning of hydrogen in oxygen is as follo'ws: 2H 2 + O 2 = 2H 2 + 116,200 cal. This means that when 4 grams of hydrogen are converted into water-vapor, 116,200 calories are liberated; 1 gram of the gas pro- duces 29,050 calories. The heat of formation of carbon dioxide is 97,000 calories, this amount of heat being produced when 12 grams of carbon burn; 1 gram of carbon furnishes, thus, 8080 calories. It is evident from this that weight for weight hydrogen furnishes over three and one-half times as much heat as carbon; for this and other reasons the gas is used as a source of heat. Hydrogen is a constituent of illuminating gas, which is used as a source of power in gas engines. The large amount of heat pro- duced when the mixture of hydrogen and air explodes raises the temperature of the gases to a high point, and as a result the pres- sure on the piston of the engine is great. The temperature of the hydrogen flame is about 2500, whereas that of burning carbon is about 1400; this difference is in part due to the fact that so much more heat is given off in the first case, and as it is formed rapidly it raises the temperature of the product of combustion to a higher point than is attained when carbon burns. This fact leads to some of the uses of hydrogen which have already been described (56). 156. It will be recalled that the oxides of certain metals, such as mercury and silver, are decomposed more or less readily when heated, and oxygen is set free; whereas other oxides such as those of sodium and calcium can be heated to the highest temperatures attainable without decomposition. A comparison of the heats of formation of these oxides will help in interpreting the facts. The values in calories are as foUows: CaO, 145,000; Na 2 O, 99,760; HgO, 22,000; Ag 2 O, 5900. In each case the values are those of the heat set free when 1 atomic weight of oxygen, 16 grams, unites with the several metals. When the oxides of the active metals are formed large amounts of chemical energy are changed into heat, and the oxides which result are very stable. In most cases if the change in energy when a compound is formed is large, the compound is a stable one toward heat, and the reverse is true. It will also be recalled that, in general, the chloride of an ele- THE ENERGY FACTOR IN CHEMICAL CHANGE 143 ment is not so readily decomposed by heat as the oxide; for exam- ple, silver chloride does not readily break down into silver and chlorine, whereas silver oxide is decomposed into its elements at a comparatively low temperature. We can make oxygen by heating mercuric oxide, but cannot use the chloride to make chlorine. On the other hand, platinum and gold, which are very inactive metals, yield chlorides which are decomposed by heat. The difference in stability toward heat between the chlorides and oxides can be traced to the fact that chlorine is a more active element than oxy- gen, and that when it unites with other elements, more chemical energy is transformed into heat than is the case in the formation of oxides. When 1 atomic weight of mercury unites with oxygen 22,000 calories are set free; when the same weight of mercury unites with chlorine to form mercuric chloride, HgC^, 54,490 cal- ories are liberated. In the case of Ag20 and 2AgCl the values are 5900 and 58,760, respectively. The heat of formation of gold chloride, AuCls, is 22,820 calories. Although the heat of formation of a compound cannot be taken as a quantitative measure of the activity of the elements composing the compound, it furnishes valuable information in many cases as to the relative activity of elements when similar compounds are compared. It should be noted that certain endo- thermic compounds are very stable toward heat. Nitric oxide, for example, is formed with the absorption of a large amount of energy, yet it resists to a high degree decomposition by heat. The relation which exists between* chemical energy and heat is considered in detail in physical chemistry and cannot be discussed further with profit at this point. CHAPTER XIII OZONE AND HYDROGEN PEROXIDE 157. The fact that the amount of energy combined with a chemical element is an important factor in determining its physical properties has been brought out in the case of carbon, which can exist as charcoal or as diamond. In the case of other elements the differences produced in this way are even more marked. Oxygen furnishes an excellent example; it exists as oxygen gas, the prop- erties of which are familiar, and as ozone, a substance that differs markedly from oxygen in chemical as well as physical properties. Ozone is a blue gas, has a marked odor, and is an exceedingly active substance; silver is not affected by oxygen, but is blackened by ozone. Ozone has a number of important and interesting applica- tions, which are based on the fact that it contains a large amount of chemical energy which can be readily transformed and, thus, utilized. The fact that a characteristic odor was formed near machines which produced electricity by friction was observed over a hundred years ago. In 1840 Schonbein showed that a distinct substance was formed, which could be made in other ways. The substance was called ozone, the name being derived from the Greek word meaning to smell. In 1856 Andrews demonstrated the fact that ozone was a new variety of oxygen. Ozone is formed when an electric discharge takes place in the air; it is formed in this way during thunder storms and near elec- trical machinery. It was thought at one time that ozone is always present in the atmosphere, but recent investigations indicate that this conclusion was arrived at as the result of using tests for the gas which were not characteristic. Ozone behaves with certain reagents like hydrogen peroxide and certain oxides of nitrogen, substances which are present in the air in small proportions. The formula assigned to ozone is Os, whereas that of oxygen is O2. The reason for this will be given later. 144 OZONE AND HYDROGEN PEROXIDE 145 Ojone y Binding Posts Connecting with Induction Coil. Oxygen orAirlnfet f/7 Oufar Tube. on Inner Tube 158. Preparation of Ozone. The gas can be conveniently prepared in the laboratory by passing air through a so-called ozonizer which is represented in Fig. 17. It consists of two con- centric glass tubes of the shape indicated. The outside of the outer tube is covered with tin-foil which is connected with one binding-post; the inside of the inner tube is covered in the same way and connected with the other. These binding-posts are joined to the terminals of an induction coil, and oxygen is forced through the space between the two tubes. As the electrical impulses flow from the metal on one tube to that on the other, they pass through the oxygen and a part of it (6 to 7 per cent) is changed into ozone. Ozonizers have been con- structed to prepare the gas on the large scale for commercial use; with these air is used, and as much as 18 per cent of the oxygen contained in it is converted into ozone. The gas is also formed when a current of elec- tricity is passed through dilute sul- phuric acid; the percentage of the ozone in the oxygen generated at the anode can be varied by changing the strength of the current and the size of the electrodes. When ozone is formed in the ways indicated above, electrical energy is transformed into chemical energy, which is taken up by the oxygen. As the result of increasing the energy associated with the element oxygen, it is transformed into a new substance possessing properties different from those of oxygen gas. The energy required to change oxygen to ozone has been determined experimentally. When this is expressed as heat energy we can represent it by the following equation: 3O 2 + 68,200 cal. <= 2O 3 When ozone decomposes into oxygen this energy is set free; as a consequence, ozone is, as we might expect, a more active oxidizing agent than free oxygen. FIG. 17. 146 INORGANIC CHEMISTRY FOR COLLEGES Ozone is formed in small quantities when potassium chlorate is heated. It will be recalled that oxygen is prepared in this way. From the standpoint of atoms it is possible that the changes involved are represented by the following equations: KC1O 3 = KC1 + 3O 2O = O 2 3O = O 3 Oxygen atoms are first liberated and then unite to form oxygen gas and ozone. At the temperature at which the reaction occurs, oxygen is much more stable, and consequently is formed in the larger quantity. That ozone is present in the oxygen generated by heating potassium chlorate can be shown by exposing to it a piece of paper on which has been put a solution of starch containing potassium iodide; ozone liberates iodine, which produces a blue color with starch. Ozone is also formed when certain elements, such as phos- phorus and zinc, oxidize slowly in moist air. It is probable that molecules of oxygen, O2, unite directly with these elements and that the resulting oxides are unstable and break down, giving the ordinary oxides and oxygen atoms, some of which unite to form ozone. 159. Physical Properties of Ozone. Ozone is a pale-blue gas which has a characteristic odor. It can be obtained as a blue liquid, which boils at 119, by passing a mixture of ozone and oxygen through a vessel surrounded by liquid oxygen. It is much more soluble in water than is oxygen; 100 volumes of water dis- solve at 0, 50 volumes of ozone and but 4 volumes of oxygen. One liter of ozone at and 760 mm. pressure weighs 2.144 grams. 160. Chemical Conduct of Ozone. Ozone decomposes slowly at ordinary temperatures into oxygen; in the presence of finely divided platinum and at about 250 the change is very rapid. It is unsafe to keep liquid ozone because it may decompose with explosive violence. Ozone is absorbed by turpentine and other oils and forms com- pounds with them; it changes rubber into a substance which is inelastic. OZONE AND HYDROGEN PEROXIDE 147 The activity of ozone has been noted; it converts silver into silver peroxide, which is black in color: 2O 3 + 2Ag = Ag 2 O 2 + 2O 2 The reaction can be used as a test for ozone provided no sulphur compound which blackens silver is present. Ozone liberates iodine from potassium iodide: O 3 + 2KI + H 2 O = 2KOH + I 2 + O 2 A solution containing potassium iodide and starch, which gives a blue color with free iodine, is used in showing the presence of ozone, but as many active oxidizing agents act in this way the reaction cannot be used as a test for the gas. From an examina- tion of the last two equations it will be seen that when ozone oxi- dizes substances, but one atom of the oxygen in each molecule of the gas unites with the substance oxidized, and that oxygen gas is formed from the remaining two atoms. This usually occurs when ozone acts as an oxidizing agent. The manner in which the atoms are linked together in oxygen and in ozone is an interesting subject for speculation. The graphic formulas usually assigned to these substances are as follows : 0=0 These formulas indicate that each oxygen atom has the valence 2. Another formula for ozone, which expresses, perhaps, its prop- erties more adequately is this: O=O=O. We can understand how a molecule made up in this way could break down into a molecule of oxygen O 2 , and an oxygen atom which would oxidize other substances; such a decomposition would be represented thus: O=O=O = O=O + The formula indicates that one of the oxygen atoms has the valence 4, a view for which there is much independent evidence. 161. Allotropic Modifications of Elements. Oxygen gas and ozone are said to be allotropic modifications of the element oxygen; 148 INORGANIC CHEMISTRY FOR COLLEGES the term is used to express the fact that the same chemical element exists as two distinct substances; they have different physical and chemical properties, and equal weights of the two possess different amounts of chemical energy. These differences are the result of the fact that they possess different atomic structures. Many elements exist in two or more allo tropic modifications; the case of carbon in the form of diamond and charcoal has already been noted. Gaseous oxygen and liquid oxygen are not allotropic forms; the meaning of the word is limited to those cases where the two substances exist in the same physical state; the two forms must be both gases, both liquids, or both solids. 162. Uses of Ozone. The fact that ozone is an active oxidizing agent has led to many applications of the gas. It will be recalled that bleaching, in many cases, is the result of oxidation. When hypochlorous acid, HC1O, is used for this purpose, hydrochloric acid is left behind in the fabric bleached. This prevents in certain cases the use of this cheap and efficient agent. When ozone bleaches, the by-product is oxygen, a gas which escapes and is not deleterious; for this reason it is of particular value in bleaching substances which are affected by active chemical reagents. It has been found of great value in bleaching flour, starch, oils, waxes, ivory, wool, and silk. Like other active oxidizing agents, ozone destroys bacteria. Water is sterilized for household use by allowing it to come into contact with air that has been passed through an ozonizer. This method is used to some extent in purifying the water supplies of several European cities; in America, bleaching powder or chlorine have found a limited use for this purpose, being preferred to ozone on account of their lower cost. Ozone is also used to purify and deodorize the air of rooms in which large numbers of people congregate. A number of theaters in Germany are equipped with apparatus to introduce ozonized air into the ventilating system. It has not yet been definitely estab- lished whether ozone destroys the bacteria and substances pos- sessing disagreeable odors that are found in the air of a room filled with people; it is possible that the ozone merely affects the sense of smell to such an extent that the presence of the unpleasant odors is not observed. The fact has been apparently established that ozone acts as an oxidizing agent only in the presence of appre- ciable amounts of water-vapor. OZONE AND HYDROGEN PEROXIDE 149 Ozone is said to have a marked effect on tobacco smoke sus- pended in the air, and its use in ventilation for clearing away the substances in smoke which affect the eyes and throat has been suggested. HYDROGEN PEROXIDE 163. In 1818 Thenard, when studying the action of acids on the oxides of metals, discovered that certain of these yielded a com- pound of hydrogen and oxygen which possessed remarkable prop- erties; the substance was shown to be a liquid and to have the composition represented by the formula H 2 O 2 ; it is now called hydrogen peroxide. Most oxides of metals react with acids to form water and salts; the reaction between barium oxide and hydrochloric acid is typical: BaO + 2HC1 - H 2 O + BaCl 2 When barium peroxide is used, however, hydrogen peroxide, and not water, is formed : BaO 2 + 2HC1 = H 2 O 2 + BaCl 2 A study of the dioxides of the metals has led to the conclusion that they can be divided into two classes: those that yield hydrogen peroxide when treated with acids, and those that do not. Exam- ples of the latter class have already been met with. When man- ganese dioxide reacts with hydrochloric acid, the products are water and manganese tetrachloride, which breaks down into chlorine and manganese chloride: MnO 2 + 4HC1 = MnCl 4 + 2H 2 O MnCU = MnCl 2 + C1 2 Manganese in MnO 2 has the valence 4 and is first converted into a chloride in which the valence of the metal is 4; at room tem- perature 2 atoms of chlorine break off from the compound, and in 150 INORGANIC CHEMISTRY FOR COLLEGES the resulting compound the metal has the valence 2. The graphic formulas of these compounds are represented as follows: y> Mn/ N) cl Mn All metallic oxides which contain 2 oxygen atoms and behave in this way are called dioxides. Oxides which yield hydrogen peroxide when treated with an acid are called peroxides. The formula of barium peroxide, BaCb, leads to the conclusion that in it the valence of barium is 4; but this is not the case. The difference in the behavior of barium peroxide and manganese dioxide is explained on the assumption that in the former compound the valence of barium is 2. The oxygen atoms are joined to the barium as indi- cated by one of the following formulas: or a== If this view is correct, the formula of hydrogen peroxide is H H or H O H The formulas of the peroxides involving an oxygen atom with the valence 4 represent better the chemical behavior of these sub- stances. 164. Occurrence of Hydrogen Peroxide. Minute quantities of hydrogen peroxide have been reported as being present in rain and snow. It is possible that the substance is formed from water and the oxygen in the air in the presence of sunlight; the bleaching of moist linen when exposed to the sunlight has been explained on the hypothesis that the hydrogen peroxide formed in this way oxidizes the coloring matters in the unbleached cloth. Enough hydrogen peroxide to be detected is formed when certain metals, such as zinc and lead, oxidize in moist air. It is probable that when substances are oxidized by gaseous oxygen, molecules of the latter OZONE AND HYDROGEN PEROXIDE 151 first add directly to them, and peroxides, which are more or less stable, are formed. In the case of sodium, the peroxide Na 2 O 2 is stable and is formed when the substance is burned in oxygen. It is probable that the peroxides of most of the metals, being unstable, break down and yield the ordinary oxide and traces of ozone or hydrogen peroxide. 165. Preparation of Hydrogen Peroxide. Hydrogen peroxide is prepared by stirring finely powdered barium peroxide into a dilute solution of sulphuric acid, which is kept cold with ice: BaO 2 + H 2 SO 4 = H 2 2 + BaS0 4 The product is filtered and diluted with enough water to make it a 3.5 per cent solution. Hydrogen peroxide decomposes more or less rapidly in the presence of bases and salts, and for this reason a trace of acid is often left in the solution. Certain organic sub- stances are thought to act as so-called negative catalyzers, that is, they retard decomposition; acetanilide is used for this purpose in certain commercial brands of hydrogen peroxide. Sulphuric acid is sometimes replaced by hydrochloric acid or phosphoric acid in preparing hydrogen peroxide; the latter yields a solution which is supposed to keep better than that obtained when sulphuric acid is used. Hydrogen peroxide is also made commercially from ammonium persulphate which is prepared by the electrolysis of ammonium sulphate (311). Hydrogen peroxide is formed when hydrogen burns in air or oxygen, although at the high temperature at which the reaction takes place it decomposes completely unless special precautions are taken. If a jet of burning hydrogen is held in contact with ice for a short time and the water formed is examined, it will be found to contain enough hydrogen peroxide to give a dis- tinct test. It is probable that a molecule of hydrogen and one of oxygen unite directly to form hydrogen peroxide: H 2 + 2 = H 2 2 It is probable, as has been stated, that when oxygen reacts with another substance the first action consists in direct combination of the two; in most cases, however, the peroxides formed are un- stable and decompose a fact which is illustrated in the case of the union of hydrogen and oxygen. 152 INORGANIC CHEMISTRY FOR COLLEGES 166. Physical Properties of Hydrogen Peroxide. Hydrogen peroxide is a syrupy liquid, which has the density 1.46. It boils at 69 when the pressure is reduced to 26 mm., and can be frozen to a crystalline solid which melts at 2. The substance is seldom isolated on account of the fact that it is apt to decompose with explosive violence. It is usually kept in solution in water. The solution ordinarily sold contains 3 per cent of hydrogen peroxide; its strength is often expressed in " volumes "; a 10- volume solution will yield ten times its volume of oxygen when decomposed. By evaporating a 3 per cent solution of hydrogen peroxide at 70 the strength can be increased to 45 per cent without much loss. 167. Chemical Conduct of Hydrogen Peroxide. When hydro- gen peroxide decomposes into water and oxygen a large amount of energy is set free : 2H 2 O 2 = 2H 2 O + O 2 + 46,200 cal. This fact leads to the conclusion that hydrogen peroxide is an active oxidizing agent. The decomposition takes place slowly at ordinary temperatures, the rate being markedly affected by rise in temperature and the presence of catalyzers such as charcoal, man- ganese dioxide, and finely divided platinum, silver or gold. Hydro- gen peroxide liberates iodine from potassium iodide; the reaction recalls the behavior of ozone with this substance : H 2 O 2 + 2KI = 2KOH + I 2 A solution containing potassium iodide and starch is used for test- ing for hydrogen peroxide; the formation of a blue color is not definite proof of the presence of hydrogen peroxide, however, since ozone produces the same effect. Hydrogen peroxide does not blacken silver, but ozone does. Hydrogen peroxide acts with certain bases as an acid; when a solution of barium hydroxide is treated with it barium peroxide and water are formed : Ba(OH) 2 -f H 2 O 2 = BaO 2 + 2H 2 O The precipitate formed is a hydrate of the composition BaO 2 ,8H 2 O. The hydroxides of certain other elements act in the same way. Some of the hydrated peroxides made in this way are used in OZONE AND HYDROGEN PEROXIDE 153 pharmacy; zinc peroxide, which is a constituent of salves, is believed to have antiseptic properties. It is a remarkable fact that with certain substances hydrogen peroxide acts as a reducing agent; for example, silver oxide is reduced to metallic silver: Ag 2 O + H 2 O 2 = 2Ag + H 2 O + O 2 It is probable that a peroxide is first formed which, being unstable, breaks down spontaneously into silver and oxygen. Certain other substances which contain a high percentage of oxygen behave in the same way. When a solution of potassium dichromate is added to one of hydrogen peroxide containing a little sulphuric acid, a blue color is developed, which soon disappears. If the solution is immediately shaken with a small amount of ether, the blue sub- stance dissolves in the latter and persists for a longer time. This behavior of hydrogen peroxide is used as a test for it. It is prob- able that the unstable blue substance formed is a perchromic acid which contains a high percentage of oxygen. 168. Uses of Hydrogen Peroxide. The uses to which hydrogen peroxide is put are based upon the fact that it is an oxidizing agent; since only water is formed from it as a by-product, it is of particular value in bleaching substances which cannot be treated with hypo- chlorites. It is used as a bleaching agent for materials of an animal source, such as wool, hair, feathers, silk, bone, and ivory. It was formerly much used as a hair bleach for toilet use. In bleaching with hydrogen peroxide the latter is not often isolated. A solution is prepared by adding sodium peroxide to a dilute solution of an acid which is kept cold: Na 2 O 2 + 2HC1 = 2NaCl + H 2 O 2 Hydrogen peroxide is much used in the household as an anti- septic and disinfectant on account of the fact that oxygen is liber- ated when it is brought into contact with blood and decaying organic matter; it is doubtful, however, whether it is of much value for this purpose. 154 INORGANIC CHEMISTRY FOR COLLEGES EXERCISES 1. What volume of ozone could be obtained from 100 c.c. of oxygen if the conversion were complete? 2. Ten liters of air were passed through an ozonizer and 10 per cent of the oxygen was converted into ozone. What was the total volume of the gases obtained? 3. Calculate approximately the volume of oxygen which is set free when 100 grams of a 3 per cent solution of hydrogen peroxide decomposes into water and oxygen. How could you express the "strength" of the solution in volumes? 4. State two ways of obtaining from a mixture of oxygen and ozone con- taining 10 per cent of the latter, a mixture of the two gases which contained a higher percentage of ozone. 5. Calculate from the formula of ozone the weight of 1 liter of the gas at and 760 mm. 6. Calculate the percentage of ozone in a mixture of oxygen and ozone 1 liter of which weighs 1.572 grams at and 760 mm. CHAPTER XIV PROPERTIES OF GASES, LIQUIDS, AND SOLIDS 169. We have seen that a study of the changes in energy which take place in chemical reactions adds materially to our knowledge of matter. Energy produces also what we have called physical changes. These are of great importance and must be treated in some detail, although they are rightly a part of physics. Physics and chemistry are so closely interwoven that the fundamentals of one science must be understood to appreciate the other; we make use of physical changes in matter to recognize and interpret chemical phenomena. The physical effect of heat energy on matter is striking; we know that we can change water into steam a gas and into ice a solid. Such changes as these are used in studying matter from the standpoint of chemistry, and in formulating theories as to the composition of molecules. We must, accordingly, have a definite knowledge of the physical properties of matter if we are to use them in helping to solve our chemical problems. 170. Gases. The characteristic property of a gas is that of diffusion; a gas distributes itself uniformly in any space into which it is put; its volume is determined by the volume of the vessel which contains it. We make use of the fact that a charac- teristic property of a gas is diffusion when we use ozone to purify the air in a room. It is only necessary to liberate some of the gas at one point, because it diffuses and finally becomes uniformly distributed. Gases are very compressible, that is, a relatively small change in pressure on a gas has a marked effect on its volume. The extent to which they expand when heated is great compared with that of liquids and solids. These facts have led to a conception of the structure of gases which has been valuable in science. According to this theory gases are made up of small particles, called mole- 155 156 INORGANIC CHEMISTRY FOR COLLEGES cules, which are in constant motion. The actual space occupied by the molecules is small as compared with the space occupied by the gas. This is evident from the fact that the volume of a sub- tance in the form of a gas is much greater than its volume when in the form of a liquid. When water is changed into steam, at atmospheric pressure, the volume of the latter is over 1500 times the volume of the water from which the steam was produced. The particles in a gas move about freely and collide with one another and the sides of the vessel which contains them. These collisions produce the pressure which the gas exerts. There is no loss in energy in the gas as the result of these collisions. When two pieces of matter strike each other a part of the energy of motion is converted into heat. If this took place in a gas it would slowly lose energy, for the heat produced would be taken up by the surroundings, and the temperature of the gas would fall. Since this does not occur, the molecules are considered to be perfectly elastic, that is, when they collide no energy of motion is changed into heat; the particle starts off after the collision with the same energy it had before. This conception of the physical structure of a gas is known as the kinetic theory of gases, on account of the fact that it postulates moving particles. The study of the physical properties of gases has led to results that make it possible to determine the size of molecules. One gram-molecular-weight of hydrogen (2 grams which occupy 22.4 liters) contains under standard conditions 6.16 X 10 23 molecules. 1 Each molecule of hydrogen weighs 0.02388 grams. 171. Liquids. Liquids are characterized by possessing a definite volume but no fixed form; they take the shape of the vessel containing them, except when in minute quantities, as drops, they assume a form more or less spherical. In many cases one liquid will mix completely with another; in such mixtures the constituents are evenly distributed. Although one liquid may be much heavier than the other, the mixture does not separate into layers. In other cases liquids are immiscible, such as oil and water. No adequate theory of the structure of liquids has been put forward. According to the kinetic theory the molecules of a liquid are packed close together and attract one another; the attraction is 1 10" is a brief way of expressing the number made up of 1 followed by 23 ciphers. PROPERTIES OF GASES, LIQUIDS, AND SOLIDS 157 not great enough, however, to prevent their motion. There is not much free space between the molecules because liquids are very incompressible a great pressure has little effect in decreasing the volume of a liquid. If the pressure on a gas is doubled, its volume decreases one-half; the change in volume of water a liquid when the pressure is changed from 1 to 2 atmospheres is 0.05 per cent. 172. Solids. Solids possess a definite form. According to the kinetic theory the molecules are closely packed and the attraction between them is great enough to prevent free motion. It is pos- sible in a solid, therefore, to have a definite arrangement of the molecules which remains fixed; this is seen in the fact that many solids assume definite forms which are characteristic. For example, when sodium chloride separates from a solution in water it appears as cubes; the molecules in building up the solid arrange them- selves in a definite mathematical way and a crystal is formed. Other substances appear in other geometrical forms; alum, for example, crystallizes in regular octahedra, which have eight triangular faces. A great variety of solid forms can be built up by combining a number of planes which cut each other at different angles. For each substance the angles at which the surfaces meet are characteristic. The study of the forms of crystals has devel- oped into a science called crystallography, a knowledge of which is of service to the chemist. 173. Crystallography. The great number of different forms of crystals that exist are classified by considering them as built up upon axes. The crystals in the regular system are referred to three axes of equal length which cut one another at right angles. If planes are passed through the ends of the axes and perpendicular to them, a cube is formed; if they are passed so that each plane cuts three axes, the solid formed has eight faces and is called an octahedron. In the tetragonal system the axes are at right angles and two" of them are equal in length. When planes are passed through the axes in the way described above, a parallelepiped is formed in the first case and, in the second, a figure composed of two square pyramids joined at their bases. The crystals in the rhombic system are built up on three axes of unequal length, which cut one another at right angles. In the monoclinic system there are two axes at right angles and a third which intersects one of these at right angles and is inclined toward the other; the axes may vary in length. Crys- tals in the triclinic system are built up on three axes which may differ in length and in the angles at which they cut one another. In the hexagonal system there are three axes in the same plane which cut one another at an angle 158 INORGANIC CHEMISTRY FOR COLLEGES of 60, and a fourth axis which passes through their intersection and is per- pendicular to the plane in which the other axes lie. This system yields hexagonal prisms and pyramids. It is readily seen that a great variety of crystal forms can be built up in the way indicated above. Crystals often show the faces of two or more forms; thus in the regular system the crystal may occur as an octahedron, the six corners of which have been removed by passing planes through the crystal perpendicular to the axes; such a form is a combination of an octa- hedron and a cube, and the crystal has, accordingly, fourteen faces eight furnished by the octahedral form and six by the cube. Great complexity can arise in this way and the subject can be followed only with the aid of models. Many solids do not apparently possess the regular arrangement of mole- cules which is thought to be present in substances that form crystals; such solids are said to be amorphous, the word being derived from the Greek word signifying without form. Most crystalline substances have a definite melting- point, that is, they change from a solid to a liquid at a definite temperature; amorphous substances either decompose before they melt, or become slowly viscous as the temperature rises there is no temperature just above which the substance is a liquid, and just below which it is a solid. 174. Properties Common to Gases, Liquids, and Solids. The characteristic physical properties of the three states of matter, we have seen, are diffusibility for gases, indeterminate form for liquids, and definite form for solids. Seeking an explanation for this we come to the conclusion that whether a substance is a solid, liquid, or gas is determined by the mobility of its molecules. All matter is made up of molecules and has, accordingly, properties which are common to the three states of matter. Many such properties have been studied with care. Only those properties which are used by the chemist in identifying substances or in inter- preting chemical transformations will be considered at length here. Many properties are carefully studied in engineering and other applications of physics. For example, what is known as tensile strength is one of the properties which determine what metal should be used in building a bridge; the electrical conductivity of copper is the property which leads to its selection as a material from which to make telephone wires. Such properties are discussed in detail in physics. The chemist uses density as we shall see, in arriving at a system of atomic weights, and the determination of the boiling- point of a liquid is the simplest practical way to identify it. These, and other properties of like significance, must be familiar to the student of chemistry and will now be discussed. PROPERTIES OF GASES, LIQUIDS, AND SOLIDS 159 175. Density. The weight in grams of 1 c.c. of a substance is called its density. On account of the fact that the density of gases is so small, the weight of 1 liter of the gas at and 760 mm. pressure is often used for this quantity. The density of water at 4 is 1, of iron 7.84, of aluminium 2.6, of sodium chloride 2.13, etc. The numbers are of great value for from them we learn the relative heaviness of substances. Since the volumes of all substances change with change in temperature and pressure, and since density is the weight of a unit volume, the density of a substance is dif- ferent under different conditions. In recording density, there- fore, the temperature and pressure should always be given unless the value is that of the substance under the conditions accepted as standard, which are and 760 mm. The change in the density of a gas under varying conditions of temperature and pressure is of such importance that it has already been discussed at great length in Chapter VIII. 176. The term specific gravity is much used. The specific gravity of a substance is a number which expresses the relation between the weight of a substance and the weight of an equal volume of some other substance. For example, if the liquid which filled a certain flask weighed 20 grams, and the water which filled the same flask weighed 10 grams, then the specific gravity of the liquid compared with water is 20 -r- 10 = 2. The expression specific gravity is a bad one, for its value is not a specific, but a relative number. The specific gravity of solids and liquids is referred to water as the standard; for gases, hydrogen or air is used. When water at 4 is the standard, the specific gravity is equal to the density, since 1 c.c. of water at 4 weighs 1 gram. 177. Specific Heat. When heat energy is supplied to matter it is absorbed, and as a consequence the temperature of the body affected rises. The specific heat of a substance is the number of calories required to raise the temperature of 1 gram of it 1 degree centigrade. A calorie is defined as the amount of heat required to raise 1 gram of water 1 degree (15-16); the specific heat of water at 15 is 1, accordingly. The specific heats of substances vary greatly; they are all less than 1. A few values are as follows: iron, 0.112; zinc, 0.093; gold, 0.032; yellow phosphorus, 0.19. The specific heat varies with temperature and pressure, and in recording the value for a substance these conditions must be stated. 160 INORGANIC CHEMISTRY FOR COLLEGES 178. The Boiling-point. If water is left undisturbed in an open vessel it will slowly disappear, evaporate; this is due to the fact that the liquid changes into water-vapor, a gas, which mixes with the air. The word vapor is commonly applied to the gaseous form of a substance which at ordinary temperatures is a liquid; vapors are gases and their behavior is, in general, in accord with the gas laws. We have seen that the gas given off from water exerts a pressure (89) and that this increases with rise in temperature. When the pressure of the vapor from any liquid just exceeds that of the air, the liquid boils; at 760 mm. pressure water boils at 100. A liquid boils when bubbles of its vapor can exist beneath the sur- face. It is evident that this can occur only when the pressure of the vapor is at least as great as that upon the surface of the liquid. The effect of pressure on the boiling-point of a liquid is marked. For example, water boils at 60 when the pressure is 14.9 cm. of mercury and at 200 when the pressure is 1169 cm. The boiling-point of a substance is a characteristic property which is often used in identifying it. Wood alcohol boils at 66 and grain alcohol at 78; it is, thus, possible to distinguish these substances from each other by determining their boiling-points. 179. The kinetic theory of gases helps materially in picturing what happens when a liquid evaporates and boils. The molecules in the liquid are in constant motion and some of those at the sur- face have enough energy due to this motion to overcome the attraction of the molecules near them and leave the liquid. They pass into the space above, forming the vapor over the liquid, and diffuse into the air. Some of the molecules in the vapor, as the result of their constant motion, strike the liquid, are attracted by the molecules there, and become again a part of it. There is thus a constant interchange, molecules passing from the liquid to the vapor and then back again. If the vessel containing the liquid is open to the air, the molecules diffuse into it and are carried away by air currents; the liquid continues to lose molecules and finally passes completely into the gaseous condition as the result of this process, which is called evaporation. If the vessel contain- ing the liquid is closed, the molecules in the form of vapor cannot escape. When the number of molecules which pass from the liquid to the vapor is equal to the number which pass from the PROPERTIES OF GASES, LIQUIDS, AND SOLIDS 161 vapor to the liquid, the two forms are said to be in equilibrium. It is important to note that the equilibrium is dynamic, not static; that is to say, it is attained as the result of equal movement in opposite directions a molecule may be in the vapor one instant and in the liquid another. A static equilibrium would result if certain molecules passed to the vapor and remained there; at equilibrium there would be no interchange between the liquid and the gas. When the temperature of a liquid is raised, the heat absorbed increases the energy of motion the kinetic energy of the mole- cules; as a result, they reach the surface with greater force and a greater number pass into the form of vapor. This is in accord with the fact that the rate at which evaporation takes place increases with rise in temperature. Since the pressure exerted by a vapor is the result of the impact of the molecules on the walls of the vessel which contains it, increase in temperature should lead to increase in vapor-pressure, a fact which has already been noted. The pressure on a liquid exposed to air is made up of the pres- sure of its vapor and of the air. As the temperature of the liquid rises, the vapor-pressure increases as we have seen; the molecules which leave the surface increase in numbers and, finally, the space above the liquid contains only molecules of vapor. At the tem- perature at which this occurs the pressure of the vapor equals that of the air; bubbles of vapor can form in the liquid, and, as a result, the liquid boils. If more heat is now applied, the liquid passes freely into vapor, and the temperature of the liquid remains constant. The heat energy is used in overcoming the attraction between the molecules of the liquid and in producing a gas against the pressure of the atmosphere. If the pressure of a gas on a liquid is increased, more molecules press upon its surface. To overcome this pressure a greater num- ber of molecules of vapor must be produced from the liquid. This is accomplished by heating it to a higher temperature; the boiling- point, thus, increases with increase in pressure. 180. Heat of Vaporization. When a liquid is heated the tem- perature rises until the boiling-point is reached; it then stays constant as the liquid passes into vapor, although heat is supplied to it. The heat energy is used in changing the liquid to a gas, and no longer shows its presence by an increase in tempera- 162 INORGANIC CHEMISTRY FOR COLLEGES ture; for this reason the heat absorbed is said to be latent. The number of calories required to change 1 gram of a liquid to vapor is called the heat of vaporization; the value varies with the tem- perature. The heat of vaporization for water at its boiling-point is 535.9 calories and that for alcohol is 206.4 calories. An important practical application of the heat of vaporization is found in ice-machines. (Fig. 18.) Ammonia gas is condensed by pressure into a liquid that boils at 33, and the heat of vaporiza- tion of which is 330 calories. When the liquid is allowed to boil, Water Condenser ~. FIG. 18. 330 calories are absorbed by each gram of ammonia as it passes into the gaseous form. This heat is extracted from the material which surrounds the vessel in which the ammonia boils. In re- frigerating machines the evaporation of the liquid ammonia takes place in a coil surrounded by a 30 per cent solution of calcium chloride, which does not freeze at the low temperature produced. The cold solution of calcium chloride is circulated around tanks containing water, if ice is to be made, or through pipes in rooms to be used for cold storage. The ammonia is pumped from the coil into pipes where it is again liquefied by pressure, and the heat given off is taken up by cold water, which flows over the pipes. It is next returned in the liquid condition to the coil and the process PROPERTIES OF GASES, LIQUIDS, AND SOLIDS 163 repeated. The heat absorbed when the liquid ammonia boils equals the heat set free when the gas is liquefied under pressure. The conversion of a gas into a liquid is the process opposite to that of vaporization; it is called condensation. In this case heat is set free, the amount being equal to that absorbed in vaporization. This fact can be represented in the case of water thus : H 2 O (liquid) + 9648 cal. ?=> H 2 O (vapor) This expression means that 18 grams of water as liquid absorb 9648 calories and pass into 18 grams of water as vapor (steam), or, read in the opposite direction, 18 grams of steam condense to water and liberate 9648 calories. The figures refer to the change at 100. A practical application of the heat of condensation is made in heating houses with steam. In the boiler-room the energy pro- duced by the burning coal is taken up as heat of vaporization by the water, the steam passes through pipes to the place to be heated, and there sets free the heat through condensation. 181. Distillation. The combination of vaporization and con- densation is much used in purifying liquids. The process, called distillation, consists in converting the liquid to a gas and then condensing it to a liquid. An apparatus for effecting these changes is illustrated in Fig. 14, p. 97. In the flask the liquid is heated to boiling, the vapor passes into the condenser, which consists of a tube surrounded by an outer jacket through which water flows. The condensed liquid is collected in the receiver. Every pure liquid which boils without decomposition has a definite boiling- point. Since the boiling-points of different substances are dif- ferent, it is possible to separate them by distillation. When a mixture of two liquids is heated the more volatile one that is, the one with the lower boiling-point vaporizes more rapidly than the other. If the vapor formed is condensed it will be found that the part that distills first will contain a much higher per- centage of the low-boiling liquid than of the one with the higher boiling-point; whereas the part which condenses last will be richer in the latter. By subjecting the various parts of the distillate the part condensed to further distillation the two liquids can, in most cases, be separated. It will be recalled that this process is used in separating oxygen and nitrogen from liquid air (22). 164 INORGANIC CHEMISTRY FOR COLLEGES 182. Liquefaction of Gases. We have seen what happens when a liquid is in equilibrium with its vapor at its boiling-point; the number of molecules that pass from the liquid to the vapor equals the number which pass from the vapor to the liquid ; vaporization and condensation are both taking place simultaneously and to the same extent. The slightest change in temperature or pressure will bring about a marked effect on the equilibrium. If we attempt to raise the temperature of the liquid by supplying heat to it and keep the pressure constant, heat will be absorbed and the evaporation will take place more rapidly than condensation; as a result the liquid will all change to vapor. If we attempt to lower the tem- perature of the liquid the evaporation will take place more slowly than condensation, and all the vapor will finally be converted into liquid. Change in pressure also brings about evaporation if the tem- perature is kept constant. If the pressure is reduced complete vaporization takes place; if it is increased the vapor changes to a liquid. The changes of conditions at the boiling-point which determine whether vaporization or liquefaction take place are exceedingly small so small that the boiling-point and liquefaction- point are practically the same. With these facts in mind the prin- ciples underlying the liquefaction of gases by pressure can be understood. Take the case of chlorine. Under a pressure of 1 atmosphere it boils at 33. If we confine some of the gas in a cylinder provided with a movable piston and by means of the latter compress it, a point will be finally reached when the gas turns to a liquid. As the pressure increases the boiling-point (liquefaction- point) rises from 33; when this reaches the temperature of the gas in the apparatus, any increased pressure causes the gas to liquefy in the way explained above. If the temperature is 20 liquefaction takes place at 6.62 atmospheres. In Faraday's original experiment on liquefying chlorine (124), the pressure was obtained by heating the solid hydrate in a closed vessel; as the compound decomposed, more and more of the gas came off and the pressure increased. 183. Critical Temperature. We have seen that the boiling- point of a liquid depends upon the pressure exerted upon it. A study of the effect of high pressures on the boiling-points of liquids has led to the discovery of a very important fact, PROPERTIES OF GASES, LIQUIDS, AND SOLIDS 165 which will be clear from the following considerations. The boiling-point of water rises with increased pressure up to 195 atmospheres, where it boils at 358. When pressures greater than 195 atmospheres are applied, water changes into a gas at 358, however great the pressure may be. Above this tem- perature water can exist only as a gas. The boiling-points of all liquids change with pressure, and for each there is a temperature above which they can exist as a gas only; this temperature is a characteristic property of liquids and is called the critical tempera- ture. The pressure which a liquid exerts at its critical temperature is its critical pressure. These constants for water are, as has been indicated above, 195 atmospheres and 358. The critical tem- peratures for a few substances are as follows: Sulphur dioxide 156, carbon dioxide 31.4, chlorine 146, oxygen 118, hydrogen 234. It will be seen that the first three gases can exist as liquids at ordinary temperatures, but that oxygen remains as a gas, what- ever the pressure upon it, at all temperatures above 118. In his experiments Faraday found that certain gases resisted his efforts to liquefy them; he called these permanent gases. It was only when it was discovered that for each gas there was a tem- perature above which it could not exist in the liquid condition the critical temperature that Faraday's unsuccessful experiments with the so-called permanent gases could be explained. In all cases he had not cooled them to their critical temperatures, and the high pressures used did not accordingly lead to liquefaction. 184. Liquefaction of Gases Possessing Low Critical Tempera- tures. In order to liquefy a gas the critical temperature of which is very low, such as oxygen, use is made of another important principle. It is necessary to do work upon a gas to compress it; a part of the energy used is transformed into heat. This fact is familiar to one who has used a pump for inflating bicycle tires; it requires work to compress the air and the pump gets warm. A large part of the heat produced results from the compression of the gas rather than from friction of the moving parts of the pump. If the heat generated on compressing a gas escapes and the latter is then allowed to expand to its original volume and pressure, it will now contain less heat energy and its temperature will, as a result, be lower than it was at first. This phenomenon is observed when the compressed air in an automobile tire is allowed to escape through 166 INORGANIC CHEMISTRY FOR COLLEGES the valve of the tire. A gas, accordingly, increases in temperature on compression and decreases in temperature on expansion. Through the application of this principle it is possible to lower the temperature of ah- to a point where its constituents liquefy. A number of forms of apparatus have been devised for this pur- pose. In one of these, air, after being freed from moisture and carbon dioxide, is compressed to 200 atmospheres by a pump and cooled; it is then sent through a coil and allowed to expand through a small opening into a second coil surrounding the first. From the second coil, in which the pressure is kept at about 15 atmospheres, the air passes through a second opening out of the machine. When the air expands from 200 to 15 atmospheres its temperature is lowered. The gas cooled in this way in passing through the second coil, circulates around that in the first, and takes up a part of its heat. As a result, in a short time the incoming gas is colder than at first, and when it expands its temperature drops lower than before. In this way the temperature of the gas at 200 atmospheres is con- tinually lowered and, finally, when it expands the temperature drops to that at which the gas is a liquid. Liquid air obtained in this way boils at about 190. In liquefying gases the boiling-points of which are very low, another principle is used. The boiling-point of a substance is lowered by decreasing the pressure upon it. Hydrogen boils at 252.5 at 1 atmosphere; if it is boiled under diminished pressure, which can be obtained by removing the gas as quickly as it is formed, the boiling-point is reduced to such a temperature that helium, which boils at 268.5, can be liquefied. 185. Solubility of Gases in Liquids; Henry's Law. Gases differ markedly in their solubility in liquids; in general, of two gases the one that is more readily liquefied is the more soluble, provided neither forms a compound with the solvent. The solubility of a gas in a liquid increases as the pressure on the gas increases. Henry studied the effect of pressure on the sol- ubility of gases in liquids and summarized the facts in the form of a law which is known by his name. The law of Henry states that the solubility of a gas in a liquid is proportional to the pressure of the gas. This law expresses the behavior of gases only when they are under a pressure which is far removed from that at which they liquefy that is, when their volumes change with pressure PROPERTIES OF GASES, LIQUIDS, AND SOLIDS 167 in accordance with Boyle's law. The weight of carbon dioxide which dissolves in water at 15 and at 2, 3, and 4 atmospheres pressure is, respectively, very nearly 2, 3, and 4 times the weight which dissolves at 1 atmosphere pressure. 186. Sublimation. Solids, as well as liquids, change directly to a gas, and exert a vapor-pressure. In most cases the pressure of the vapor at ordinary temperatures is very small, but in many it is great enough to be recognized by the fact that the solid possesses a characteristic odor. For example, camphor, iodine, and naph- thalene, which is used in " moth-balls," all give off at room tem- perature appreciable quantities of vapor. By cooling the vapor in each of these cases it can be changed directly to the solid. The transformation of a solid to a gas and back to the original solid without passing through the liquid state, is called sublimation. A number of solids do not melt when they are heated, because the vapor-pressure of the gaseous form produced equals the pressure of the atmosphere at a temperature below the melting-point of the solid. Application is made of this principle in the purification of iodine, ammonium chloride, naphthalene, and other substances of commercial importance. 187. The Melting-point. The change of a liquid to a gas and that of a solid to a gas have been discussed above; we shall now consider the relation between the liquid and the solid state. The temperature at which a crystalline solid changes to a liquid is called its melting-point, and that at which a liquid changes to a solid is called its freezing-point. For any one substance these points are the same. If a solid is heated it rises in temperature until a point is reached where it begins to melt ; this temperature is called its melting-point. Any additional heat applied does not affect the temperature, but transforms more of the solid into liquid. The number of calories required to convert 1 gram of a solid into a liquid is called its heat of fusion. The heat of fusion of ice is 79 calories. If a liquid is cooled its temperature falls to the freez- ing-point, and stays constant until all the liquid has changed to solid. The heat lost in solidifying is equal in amount to that absorbed in liquefaction. Pressure has, as we have seen, a marked effect on the boiling- point of a liquid; it has but little effect, however, on the freezing- point. This is due to the fact that there is a great change in volume when a liquid is converted into a gas, and the latter in 168 INORGANIC CHEMISTRY FOR COLLEGES forming must overcome the pressure upon it. In the change from the solid to liquid state or the reverse, the volume change is very small. The boiling-point of water is raised about 20 degrees by increasing the pressure upon it from 1 to 2 atmospheres. The same increase brings about a lowering of only 0.008 degree in the freezing-point. The boiling-point of a liquid always rises with in- creased pressure, because the volume of the gas is always greater than that of the liquid. The melting-point of a substance may be raised or lowered by increasing the pressure; if the volume of the liquid is greater than that of the solid, increase in pressure raises the melting-point; if the volume of the liquid is less than that of the solid, increase in pressure lowers the melting-point. The use of ice as a refrigerating agent is based on the fact that in melting it absorbs heat, and thus cools the surrounding air. When liquid air was first made it was thought that it would be a valuable cooling agent to replace ice on account of its very low boiling-point. A consideration of the heat changes involved when liquid air boils and ice melts will bring out the fact that the former is not so good an agent as ice. When 1 gram of liquid air at 190 changes to a gas at this temperature, about 50 calories are absorbed. To raise this air to requires 18 calories. One gram of liquid air will thus absorb 68 calories in changing to air at 0. At this temperature 1 gram of ice absorbs 79 calories when it melts; it is evident, therefore, that ice is more efficient in absorbing heat than liquid air. Water is characterized by the fact that it requires such large amounts of heat to change it from one state to another. This fact makes ice an efficient refrigerating agent. 188. Transition-points. When a substance is at its melting- point the solid and liquid forms of it exist, side by side, in equilib- rium, if heat is neither added nor taken away; the solid is melting and the liquid is freezing, the two processes taking place at the same rate. If heat is now added or taken away, one form changes to the other without a change in temperature. The temperature at which this occurs is called a transition-temperature ; in the case cited it is a melting-point. The boiling-point is also a transition- point, for at this temperature a liquid changes into a gas as heat is applied without a change in temperature taking place. The word phase is often used in describing the substances which undergo changes of this kind; water can, for example, exist in three forms PROPERTIES OF GASES, LIQUIDS, AND SOLIDS 169 or phases gaseous, liquid, and solid. At the melting-point the liquid and solid phases are in equilibrium, and at the boiling-point the liquid and gaseous phases. A phase is any homogeneous constituent of a system made up of either different physical states of the same substance or of different substances. In sublimation we have two phases, the gaseous and the solid. If we had a solu- tion containing salt and some solid salt in a bottle half filled with the liquid, we would have three phases, the gaseous, the solution, and the solid. Many substances can exist in more than three phases, for example, sulphur exists as a gas, in two distinct liquid forms, and in two solid crystalline forms. The temperature at which any one of these forms or phases is in equilibrium with any other phase is a transition-point. 189. Superheating and Supercooling. If water is carefully heated in a polished dish of silver it is possible to raise its tempera- ture above 100 without boiling taking place. Globules of water suspended in oil have been heated to 145 without passing into steam. The boiling-point of a liquid is not the highest tempera- ture to which it can be heated, but the temperature at which the vapor and the liquid are in equilibrium. Water at a higher tem- perature than this is said to be superheated. If a trace of the gaseous phase is introduced into superheated water, the change to steam takes place at times with explosive violence, and the tem- perature drops to 100. It is evident, therefore, that water can- not be superheated in the presence of steam; and if it is desired to raise it above 100 care must be taken to prevent a bubble from forming. This is accomplished by using a polished vessel free from inequalities upon which a bubble can form. In order to prevent superheating and " bumping " when boiling a liquid in a glass flask, pieces of broken porcelain or glass are introduced to furnish sharp points upon which bubbles of steam may form. Liquids can also be supercooled that is, kept in the liquid condition below their freezing-points. If a bit of the solid phase is introduced into the liquid the mass quickly freezes; the same effect is produced by agitation or by the introduction of a rough surface upon which a crystal of the solid can form. 170 INORGANIC CHEMISTRY FOR COLLEGES EXERCISES 1. Upon what physical properties of the substances involved are based the use of the following: (a) mercury in thermometers, (6) a double boiler in cooking, (c) sulphur dioxide as a disinfectant, (d) steel in bridge construc- tion, (e) heating of houses with hot water, (/) wooden handles on soldering irons. 2. Name two conditions when it would be advisable to use a gas in a thermometer instead of mercury; give a reason in each case. 3. Why is it more economical to run an engine with steam at high pres- sure than at low pressure? 4. Name in each case one use of a substance which is based in part on the following: (a) high density, (6) low density, (c) high specific heat, (d] high latent heat of fusion. 5. If a piece of iron at room temperature is put into 100 c.c. of hot water what would be the effect on the temperature of the water? Why? If the experiment were repeated using a piece of gold of the same weight and tem- perature would the quantitative result be the same? If not, why? 6. What difference would you observe if you drank hot tea in one case from a cup made of porcelain and in the other from a cup made of alum- inium? Why? 7. Calculate the weight in grams of cubes of the following substances, each 60 cm. long on the edge: copper, lead, tin. The densities of the metals, are, respectively, 8.95, 11.34, and 7.3. 8. A piece of wood which measured 2 X 4 X 12 inches was found to weigh 2 Ibs. Calculate- the specific gravity of the wood. 9. A piece of apparatus constructed of steel was found to weigh 750 Ibs. What would it weigh if it were constructed of aluminium and its dimensions were the same? The specific gravities of steel and aluminium are 7.6 and 2.7, respectively. 10. If equal weights of alcohol and water are allowed to stand in open ves- sels which liquid would disappear first? Why? 11. Certain substances which are decomposed when an attempt is made to distill them in the type of apparatus represented in Fig. 14, can be dis- tilled unchanged if the apparatus is freed from air before the process is carried out. Can you explain this fact? 12. If a mixture of water and aniline, which is a liquid that boils at 182, is placed in the distilling apparatus represented in Fig. 14 and heated to boiling, and the vapor is condensed, the product is a mixture of the two liquids. Explain why aniline which boils at 182 can be "distilled with steam" in this way. 13. Would there be any difference in the cost of distilling 1000 gallons of water and the same volume of alcohol? If so, why? 14. How could you prepare relatively pure carbon dioxide from the mix- ture of gases obtained when carbon is burned in air, by using a process based on Henry's law? CHAPTER XV CARBON AND ITS OXIDES 190. Carbon is one of the most important and interesting ele- ments, because its compounds are so widely distributed in nature and play such an important part in daily life. All living things contain carbon, and life itself consists of a series of chemical changes in which carbon compounds take part. In organic chemistry the compounds of this element are studied in detail; over 200,000 of these are known and many have been carefully investigated. The transformations which chemists have brought about in carbon compounds, have led to the preparation of thousands of sub- stances which have been put to valuable uses; these include dyes of many hues, medicines, antiseptics, photographic developers, soaps, edible oils, inks, perfumes, and many other substances which add to our comfort and pleasure in living. Through a study of carbon compounds we learn what takes place in the digestion and assimilation of foods; we discover the difference in value of bread and meat in supporting life, and how much of the various kinds of food a man must eat to remain strong and be able to work. The more the chemistry of the compounds of carbon is studied the nearer we get to an understanding of life itself the greatest of the unsolved problems of science. Before such interesting subjects can be considered, however, it is necessary to study carbon itself and the simpler compounds of the element. Carbon compounds are widely distributed in the inorganic world. Great mountain ranges are made up of calcium carbonate and mag- nesium carbonate, and carbonates of other metals are important ores from which these metals are extracted for industrial purposes. Natural gas and petroleum are made up of compounds of carbon, and coal from which we derive the energy to do the world's mechan- ical work, is a mixture of carbon and some of its derivatives. 171 " 172 INORGANIC CHEMISTRY FOR COLLEGES 191. Diamond. Carbon occurs in the free condition as dia- mond, the chief sources being South Africa, Brazil, and India. Diamonds are found as crystals, but the crystalline form is more or less indistinct, and they generally resemble rough pebbles. When perfectly pure they are colorless, although some that are slightly colored by impurities, blue, green, or yellow, are used as gems. Black diamonds are used for grinding purposes. In order to bring out the brilliancy of the diamond it must be cut and pol- ished. Faces are ground upon it at such angles that the maximum amount of the light which penetrates the diamond is reflected from the rear faces back through the stone. This light in passing into the air is broken up into the various colors of the rainbow. A " brilliant " is a stone ground in the shape of a many-sided pyramid around the edge of the base of which are cut faces; the flat surface is the face of the diamond. Since the diamond is the hardest substance known it must be polished with diamond dust. Many diamonds have become famous on account of their size, or the historical personages who owned them. The largest dia- mond known is the so-called Cullinan; it was found in South Africa and weighed 3032 carats, which is almost 1^ pounds. This priceless gem was presented to King Edward VII of England. It was cut into a number of jewels, the largest of which weighs 516 carats. The famous Kchinoor, which is one of the crown jewels of England and belonged to Queen Victoria, weighs 106 carats. The international carat, which weighs 200 milligrams, has now replaced the older English carat of 4 grains (205 mg.). The cost per carat of a diamond increases rapidly with the size. Diamond is insoluble in all liquids. It is stable at ordinary temperatures, but when heated out of contact with the air, changes to graphite. At high temperatures it burns in oxygen, and carbon dioxide is formed. Diamond has the density 3.5, and does not conduct electricity. The preparation of diamonds artificially has always been an interesting problem to the chemist. Moissan, a French chemist, succeeded in 1887 in making diamonds, but the process did not yield a product of commercial value. Moissan's experi- ments were of great scientific interest, however, for they dem- onstrated how one allotropic form of carbon could be changed CARBON AND ITS OXIDES 173 to another, and showed the relationship between amorphous char- coal and crystalline diamond. Charcoal dissolves in about 20 times it weight of melted iron. When the solution cools a part of the carbon separates out in crystals as graphite, which is a third form of carbon. The density of graphite is 2.3, which is consider- ably less than that of the diamond (3.5). It appeared possible to have the carbon separate in the denser form from its solution, if great pressure were exerted upon it during crystallization. When a mass of molten iron cools it solidifies first on the outside, and a solid shell is formed which increases in thickness as the solidifica- tion takes place. Iron expands on passing from the liquid to the solid condition, and, as a consequence, when the liquid core of a cooling mass of iron solidifies, it exerts a tremendous pressure on anything in it, since the solid shell prevents expansion outward. Moissan found that when carbon separated from melted iron which was suddenly cooled, it appeared, in part, as diamonds. These were obtained by dissolving the iron in acid, and were found to be very small, the largest having a diameter of about one-half a millimeter. It is of interest to note here that meterorites com- posed of metallic iron have been found which contained graphite and diamonds, both black and colorless. The way to prepare diamonds commercially will no doubt be discovered. Artificial rubies, sapphires, and other gems have been made, which equal in beauty those found in nature. Diamonds which are of no value as jewels on account of their color, are used to cut glass, as tips for drills to bore rocks, and as a powder to polish jewels. The dark-colored or black form of the diamond, which may occur in pieces weighing half a pound, is called carbonado or bort. 192. Graphite. Graphite is a more or less pure form of car- bon, which occurs in crystalline masses. It is black or steel gray in color, and is made up of flakes which are soft and slippery. Its specific gravity varies, that of the purest varieties being about 2.3. Graphite is widely distributed in small quantities; it is mined for commercial purposes in New York, Canada, Siberia, and Ceylon. The uses of graphite are based upon its physical properties and its chemical inertness. It changes only slowly when heated in the air, and for this reason is used for making crucibles and other apparatus in which chemical reactions are to 174 INORGANIC CHEMISTRY FOR COLLEGES be carried out at a high temperature. Graphite conducts elec- tricity and is but slowly attacked by chlorine; these properties make it available for electrodes in the electrolytic industries. The name given to graphite was derived from the Greek word meaning to write, on account of the fact that due to its softness it left a black mark when drawn across paper. It was also called plumbago or black lead, because lead possesses this property. In making lead pencils ground graphite is mixed with clay and forced through dies; the sticks are then heated to a high tempera- ture. The hardness of a pencil is determined by the proportion of graphite mixed with the inert material. Graphite is used in making stove polish, because it protects the iron of the stove from the action of oxygen, and even at high temperatures changes but slowly; and owing to its crystalline structure it reflects light and thus serves as a polish. The use of graphite as a lubricant is based upon the fact that it is made up of slippery scales. It can be used when the temperature is high, when dust is apt to collect if oil is the lubricant, and for reducing the friction in machinery made of wood; in the last case oil would be rapidly absorbed by the wood, whereas graphite is not. On account of the fact that graphite conducts electricity, it is used to cover molds of wax which are to be coated electrolytically with metals in making electrotypes. A large amount of graphite is now manufactured from coal at Niagara Falls by a process invented by Acheson. Ground anthra- cite coal, with which is mixed some iron oxide and sand, is placed in a furnace made of fire-brick. By means of electrodes which enter the furnace through the walls, a powerful alternating cur- rent of electricity is passed through the mixture. On account of the granular structure of the material in the furnace heat is gen- erated as the result of the resistance offered to the current. Elec- tric furnaces which are heated in this way are called resistance furnaces; they are much used in the industrial preparation of compounds requiring very high temperatures for their formation. When the mixture in the furnace is heated in this way for twenty- four hours, it is changed to graphite. As the heat in the furnace is increased some of the carbon first unites with the oxygen present in the iron oxide and sand (silicon dioxide) and carbides of iron and silicon are formed; at the high temperature finally attained the iron and silicon are volatilized from these compounds and the CARBON AND ITS OXIDES 175 carbon is left as graphite. By molding the materials used with pitch into the particular forms desired, and then baking them at a high temperature, it is possible to prepare electrodes, which are changed to graphite when heated in the electric furnace. 193. Wood Charcoal. Charcoal is a more or less pure form of carbon, which is obtained by decomposing at a high temperature substances which contain compounds of carbon. The proper- ties of the resulting charcoal are determined by its source. Char- coal lacks crystalline structure; it is said to be amorphous. The purest variety is obtained by heating sugar to a high temperature away from the air. Larger quantities of charcoal are prepared from wood, on account of the cheapness of the source. Logs of wood are piled closely together and covered with turf; they are then set on fire and the amount of air admitted is such that the wood smolders. At the end of a number of days the fire dies out and the wood which has not burned is recovered as charcoal. Very valuable products are lost in this way, and with the increasing value of wood, charcoal is made more and more by heating wood in iron retorts so arranged that the volatile products formed can be con- densed and saved. Since no wood is burned the yield of charcoal is greater than when a kiln is used. Carbon is a very inert element at ordinary temperatures. Charcoal resists the action of air and moisture, and for this reason piles and the ends of fence posts which are to be placed under- ground are often charred. It is said that parts of a bridge erected by Caesar over the Thames River were discovered in modern times, and the wood was found to be charred. Charcoal is used as a fuel in the household; it is of particular value for broiling because it burns without smoke and gives a high heat. It is also used in the preparation of iron from its ores; the carbon in this case unites with the oxygen combined with the metal. Charcoal is one of the ingredients of gun-powder (620). Charcoal retains the shape of the wood from which it was prepared. Since large quantities of volatile material are lost in converting wood to charcoal, the latter contains a mass of inter- stices and channels from which these materials have been driven out. As a result, the charcoal is porous and has a very large sur- facea fact which leads to important properties of carbon in this form. When gases come into contact with solids, some of the 176 INORGANIC CHEMISTRY FOR COLLEGES former adhere more or less strongly to the surface of the solid in a way that is not as yet understood. If air is passed through a glass tube containing hydrogen we should expect that all of the latter would be driven out. If this is done at room temperature, and the gas tested from time to time, it will be found that no hydrogen can be discovered in the gas. If now the glass tube is heated to a higher temperature and the gases tested again, hydrogen will be found to be present. This process can be repeated until very high temperatures are reached, and each time increasing the tem- peratures causes more of the gas to appear. The conclusion is drawn that in some way the molecules of a solid at its surface attract the molecules of a gas and hold them by what is called adsorption. The phenomenon is one that takes place at the sur- face, and as a consequence charcoal has high adsorptive powers. The amount of gas held in this way is determined by the kind of wood from which the charcoal is made, as this is a factor in the porosity of the charcoal. The extent to which a gas is adsorbed is also determined by the nature of the gas; in general, the more readily a gas is condensed by pressure, the more it is adsorbed. Charcoal made from boxwood or cocoanut shells shows a high power of adsorption. Boxwood charcoal adsorbs 90 times its volume of ammonia, 55 volumes of hydrogen sulphide, and 9 volumes of oxygen. Adsorbed gases are more reactive than those in the usual condition. If charcoal containing chlorine is brought into hydrogen the gases unite at ordinary temperatures. If dogwood charcoal, as soon as it is made, is powdered, it takes fire as the result of the fact that the heat generated in the adsorption of the air causes the oxygen and carbon to unite; charcoal that behaves in this way is said to be pyrophoric. Powdered charcoal made into pellets is used medicinally to relieve the pain caused by the pressures of gases in the stomach and intestines resulting from indigestion. Charcoal is used to produce a very high vacuum. When it is in a vessel from which as much air as possible has been pumped out in the usual way, it adsorbs a large part of what is left. Adsorbed gases can be removed unchanged from charcoal by heating it in a vacuum to a high temperature. The high adsorbing quality of charcoal for gases was utilized in the construction of gas masks used in warfare. As the result of extended experimentation char- CARBON AND ITS OXIDES 177 coal was prepared in such a form that it possessed markedly increased adsorbing power over that which had been previously prepared. This was accomplished by heating the charcoal to a high temperature in steam, and by adding to it small quantities of other substances. Charcoal is used as a catalyzer to facilitate the union of gases; the explanation of its behavior here is traceable to its power to adsorb one or both of the gases. In the highly condensed condi- tion in which they exist when adsorbed, they react at ordinary temperatures. Such facts as these add great interest to the study of the causes underlying adsorption, and attention is now being devoted to this highly interesting and important phenomenon. When it is more fully understood we shall be nearer an explana- tion of certain kinds of catalytic action. Charcoal was used as a catalyzer in preparing phosgene, an important war-gas, which has the formula COCk. When chlorine and carbon monoxide, CO, are mixed and exposed to the sunlight, the gases combine slowly. This method was inapplicable when large quantities of phosgene had to be prepared. It was found that the passage of the gases over charcoal brought about their union readily, and this method was used. 194. Animal Charcoal. Charcoal also adsorbs solids and liquids. Coloring matters can often be removed from solutions by passing the latter over charcoal. For this purpose charcoal made by heating bones to a high temperature is very efficient. The charcoal in this case is spread over the mineral matter of which bones are chiefly made up calcium phosphate. Charcoal made in this way is called bone-black or animal charcoal. Crude sugar contains a brown substance; when it is dissolved in water and the solution is allowed to trickle over bone-black the substances which impart a color to it are completely adsorbed. From the colorless solution pure white sugar, as we know it, is obtained by evaporation and crystallization. Coloring matters like litmus, indigo, and those in vinegar, tea, etc., are readily adsorbed by charcoal. In general, substances whose molecules are large are adsorbed. The organic matter in drinking water is also adsorbed by charcoal, a fact which was formerly utilized in the household; but as the niters containing charcoal soon become clogged and ceased to act, the use for this purpose has been discontinued. 178 INORGANIC CHEMISTRY FOR COLLEGES Charcoal which is used industrially for decolorization is periodi- cally heated to a high temperature to restore its activity. 195. Lampblack. When many substances containing a large percentage of carbon are burned in a limited supply of air, the other elements present unite with the oxygen, and carbon is left in the state of a very fine powder. Amorphous carbon prepared in this way is called lampblack; it is familiar as the soot formed when a lamp smokes. Lampblack has a number of important industrial uses and is prepared from residues obtained in the purification of oils from petroleum and coal-tar. The products formed as the result of incomplete combustion are led through a series of cham- bers in which the soot deposits. Natural gas is also a source of lampblack. It is prepared by allowing the burning gas to fall on a rotating cold iron plate from which the carbon is scraped off. This so-called gas-black is used in the manufacture of some forms of black rubber for automobile tires. Lampblack is used in making paints, shoe polish, printer's ink, India ink, and other substances which require a black pigment. Coke is an impure form of carbon which is obtained by heating coal. It is of such importance that it will be treated in detail in the next chapter. 196. Relation between the Allotropic Forms of Carbon. The fact that equal weights of diamond and charcoal give different amounts of heat when burned has already been noted. One gram of diamond produces when burned 7805 calories, 1 gram of graphite 7850, and 1 gram of sugar charcoal 8040 calories. Diamond and graphite exist as crystals and have a smaller heat of combustion than charcoal; it is evident that the orderly arrangement of the atoms takes place with the evolution of heat. This arrangement leads to a closer packing together of the atoms, for the density of the diamond is 3.5, that of graphite about 2.3, and that of charcoal about 1.5. One gram of carbon as charcoal occupies over twice the space occupied by 1 gram as diamond; this does not refer to the volume of the charcoal, which includes air spaces in its pores, but only to that occupied by the carbon itself. The difference in the physical condition and energy content of the allotropic forms of carbon leads to a difference in chemical activity. Diamond is scarcely attacked by the most vigorous chemical reagents, such as powerful oxidizing agents, whereas CARBON AND ITS OXIDES 179 graphite is slowly, and charcoal more rapidly converted into sol- uble substances by a mixture of nitric acid and potassium chlorate. 197. Properties of Carbon. Many of the properties of carbon have been noted above. Carbon does not melt; at the tempera- ture of the electric arc it volatilizes, and even at lower tempera- tures sublimes slowly. The black coating that appears in incandescent electric bulbs in which a carbon filament is used, is a deposit of the element. Under ordinary conditions carbon is one of the most inert of elements, but as the temperature is raised its activity increases and, finally, at the temperature of the electric arc it is one of the most active of elements. When heated sufficiently it unites with sulphur, silicon, iron, aluminium, calcium, and other elements; the resulting compounds are called carbides. Many of these which are now of great industrial importance were first made by Moissan, who applied electricity to the production of high tem- peratures, and invented methods of studying chemical change at 3000 and over. The more important carbides will be described later. CARBON DIOXIDE 198. On account of the fact that carbon dioxide is present in the air and plays such an important part in animal and vegetable life, it is a compound of the first importance. It was one of the first gases to be recognized as being distinct from air. Van Hel- mont (1577-1644) studied the gas produced during alcoholic fer- mentation and showed that it was also produced by the action of acids on chalk and when carbon burned. He called it " gas sylvestre." In 1757, Joseph Black, a Scotchman, showed that carbon dioxide was absorbed by alkalies like sodium hydroxide, and as the gas disappeared he called it " fixed air." He demon- strated that it was obtained when limestone was heated to make lime, and that an insoluble compound was formed when it was passed into limewater, a reaction which is used to-day as a test for the gas. Priestley discovered the presence of carbon dioxide in the air, and Lavoisier showed that it was formed in breathing. 199. Occurrence of Carbon Dioxide. The gas is a normal constituent of the atmosphere being present to the extent of 3 parts by volume in 10,000. Its presence in the air is largely due 180 INORGANIC CHEMISTRY FOR COLLEGES to the fact that it is produced as the result of the decomposition of organic material, such as decaying leaves, trees, and other vege- table material; it is also produced in respiration and when carbon compounds burn. Carbon dioxide is present in large quantities in the gases which issue from volcanoes; it is given off from fis- sures in the earth's surface, and on account of the fact that it is heavier than air it often collects in caves and low-lying confined valleys. In the Grotta del Cane near Naples the bottom of the cave is said to be covered with a layer of air about 18 inches deep containing a large proportion of carbon dioxide. If a dog enters the cavern it is soon suffocated, but a man can explore it in safety. Death Valley in the Yellowstone Park is so called because animals have died there as a result of the fact that the air was rendered unfit for breathing by carbon dioxide. 200. Preparation of Carbon Dioxide. Carbon dioxide is formed when carbon and any of its compounds are burned in an ample supply of air: C + O 2 = C0 2 If the compound contains hydrogen, water is also formed, for example, the reaction which takes place when methane burns is as follows: CH 4 + 2O 2 = C0 2 + 2H 2 O Carbon dioxide is used to prepare other compounds of commercial importance. For this purpose the gas is prepared by heating limestone, which is calcium carbonate: CaCO 3 = CO 2 + CaO Calcium oxide, which is obtained along with carbon dioxide, is quicklime; it was formerly called caustic lime, the word caustic being derived from the Greek and signifying burnt; lime was pre- pared by " burning " limestone. Other carbonates decompose into carbon dioxide and the oxide of the metal, or the metal itself, when heated. The temperature at which decomposition takes place freely varies with the different carbonates. It will be recalled that the less active a metal, the more readily its oxide is decomposed by heat; the same is true of the carbonates although these compounds are broken down more readily than the oxides. Calcium oxide has not been decomposed by heat, but its carbonate CARBON AND ITS OXIDES 181 can be converted into the oxide of the metal and carbon dioxide; at 900 the pressure of the gas given off is equal to that of the atmosphere. The carbonates of the active metals, sodium and potassium, decompose only very slowly at the temperature of the electric furnace. When the carbonate of a very inactive metal like silver is heated, the temperature of decomposition is such that the oxide cannot exist, and, as a consequence, oxygen is also lost and the metal is left in the free condition. Carbon dioxide is also prepared by the action of acids on cal- cium carbonate; this is the method usually employed in the labora- tory: CaC0 3 + 2HC1 = CaCl 2 + H 2 C0 3 H 2 CO 3 = H 2 O + C0 2 A double decomposition first takes place as the result of which carbonic acid, H^COs, and a salt of calcium are formed. The former is unstable and breaks down into carbon dioxide and water. The reaction takes place as the result of the fact that one of the products formed escapes. All carbonates are decomposed by acids, and the formation of carbon dioxide when they are treated in this way is used as a test for carbonates. Carbon dioxide is formed when certain organic substances, like the sugar present in fruits, undergo fermentation. The gas formed in this way is collected and, stored under pressure in cylinders, is an article of trade. Carbon dioxide is formed in the putrefaction and decay of all organic substances, and most of that which gets into the air comes from this source. It is also formed in respiration. 201. Physical Properties of Carbon Dioxide. Carbon dioxide is a colorless gas, which is odorless when diluted with other gases; the pure substance, when breathed, produces a slight tingling sensation. It is about 1.5 times as heavy as air; 1 liter at 760 mm. and weighs 1.976 grams. On account of this fact, the gas can be collected by upward displacement of air, and can be poured from one vessel to another like water. Carbon dioxide dissolves in water; the amount which passes into solution is determined by the temperature of the water and the pressure exerted on the gas. At 15 and 760 mm. pressure 1 liter of water dissolves 1 liter of the gas; at 2 atmospheres, twice the weight, at 3, three times; at mod- 182 INORGANIC CHEMISTRY FOR COLLEGES erate pressures the solubility is proportional to the pressure of the gas. A solution of carbon dioxide in water under a pressure of 3 to 4 atmospheres is sold as soda water. The name is derived from the fact that baking soda, which is a carbonate, was formerly the source of the carbon dioxide used in making the water. When a bottle containing soda water is opened the pressure is reduced to that of the atmosphere, and the carbon dioxide comes off as bub- bles, until the amount in solution is reduced to that which water can dissolve at the temperature of the water and the pressure of the atmosphere. 202. As the critical temperature of carbon dioxide is 31.4, it can be liquefied by pressure alone at ordinary temperatures. Liquid carbon dioxide has the specific gravity 0.95 at 0; it has a vapor-pressure of 59 atmospheres at 20 and must be kept in cyl- inders of steel that can resist this pressure. When liquid carbon dioxide comes into the air it evaporates very rapidly and much heat is absorbed; this heat is taken up, in part, from the sub- stance itself, the temperature of which accordingly falls, and when 79 is reached it solidifies to a white, snow-like solid. It first becomes a solid rather than a liquid because at the pressure of 1 atmosphere there is no temperature at which gaseous and liquid carbon dioxide are in equilibrium; it has no boiling-point. If solid carbon dioxide at a low temperature is heated, its vapor-pressure increases until at 79 it equals the pressure of the air, and, at this point, sublimes without first changing to the liquid state. Under a pressure of 3.5 atmospheres liquid carbon dioxide boils at 56. Solid carbon dioxide is used as a means of securing low temperatures for experimental work. It is mixed with ether or acetone in order to bring the cooling agent in close contact with the vessels used to contain the materials being studied. At 79 mercury is a solid, and a bit of rubber tubing becomes so brittle it can be broken. 203. Chemical Properties of Carbon Dioxide. Carbon dioxide is a very stable compound. When heated to 2000 it dissociates to the extent of about 1.8 per cent into carbon monoxide and oxygen: 2C0 2 <^ 2CO + O 2 If the temperature is lowered the products of the dissociation unite to form carbon dioxide. When carbon dioxide is led over the CARBON AND ITS OXIDES 183 more active metals heated to a high temperature, a decomposition like that produced by heat takes place, but the oxygen unites with the metal: Zn + C0 2 = ZnO + CO Since carbon dioxide is the product formed when carbon burns, the gas does not support the combustion of carbon or its compounds. If a piece of burning wood is put into carbon dioxide it is imme- diately extinguished. Wood will not burn in air which contains as little as 15 per cent by volume of carbon dioxide. A solution of carbon dioxide in water shows the properties which are characteristic of acids; it turns litmus from blue to red, and forms salts with bases. A reaction takes place between carbon dioxide and water as the result of which carbonic acid is formed: CO 2 + H 2 O <= H 2 C0 3 The reaction is a reversible one; as the temperature is raised the carbonic acid dissociates into water and carbon dioxide, which escapes as a gas from the solution. Carbon dioxide is classed as an acid anhydride, because it is formed as the result of the elimina- tion of water from an acid. Carbonic acid reacts with sodium hydroxide to form sodium carbonate : 2NaOH + H 2 C0 3 = Na 2 C0 3 + 2H 2 O A different salt can be formed by using the amounts of the two substances represented by the following equation: NaOH + H 2 C0 3 = NaHC0 3 + H 2 O In this case but one-half of the hydrogen of the acid is replaced by sodium; the salt is called sodium bicarbonate. Acids, in general, which contain 2 hydrogen atoms can form two kinds of salts; they can react with one molecule and with two molecules of a base like sodium hydroxide; and for this reason are called dibasic acids. Hydrochloric acid, HC1, which contains but 1 hydrogen atom is a monobasic acid and phosphoric acid, H 3 PO4, which contains 3 atoms of hydrogen that can be replaced by metals, is called a tribasic acid. The reaction which takes place between carbon dioxide and sodium hydroxide serves as a convenient means of separating it 184 INORGANIC CHEMISTRY FOR COLLEGES from other gases. If, for example, a mixture of oxygen and carbon dioxide is passed through sodium hydroxide, the latter reacts and sodium carbonate, a solid soluble in the water, is formed; the oxygen passes on unchanged and in this way is freed from the car- bon dioxide. 204. Carbonates. The carbonates of many metals occur in nature, some of which are important minerals and will be described later. The carbonates belong to three classes the normal car- bonates, like Na 2 CC>3, CaCOs, and FeCOa, in which no hydrogen is present, the acid carbonates, like NaHCOa, which are formed as the result of the partial replacement of the hydrogen in carbonic acid by metallic atoms, and the basic carbonates. To the last- named class belong the salts formed as the result of the partial replacement of the hydroxyl groups of bases by the radical of carbonic acid, COs; the composition of basic salts is usually com- plex and will be considered later. All the normal carbonates of the commonly occurring elements are insoluble in water except those of sodium, potassium, and ammonium. It will be well to remember this statement, as it summarizes many facts which will be useful later. The insoluble carbonates can, in general, be made by double decomposition, as illustrated here by the case of calcium carbonate : CaCl 2 + Na 2 CO 3 = CaCO 3 + 2NaCl 206. Test for Carbon Dioxide and Carbonates. When carbon dioxide comes in contact with a solution of calcium hydroxide, Ca(OH)2, calcium carbonate, an insoluble substance, is formed: CO 2 + Ca(OH) 2 = CaCO 3 + H 2 O This reaction serves as a test for carbon dioxide, because no other gas behaves in this way. The test is made by placing a glass rod into a solution of calcium hydroxide (lime-water), and then holding the adhering drop of the solution in the gas. If carbon dioxide is present the surface of the drop will be covered with a white coating. If it is desired to determine whether a substance is a carbonate, a bit of it is treated with dilute hydrochloric acid in a test-tube, and a rod moistened with lime-water is held in the gas formed. If an appreciable quantity of the gas is set free, it can be poured after a few seconds into another tube about one-third filled with lime- CARBON AND ITS OXIDES 185 "JfS0 4 NaHCOj Solution '" water, care being taken to transfer the gas only. The tube con- taining the lime-water is closed with the thumb and shaken; if carbon dioxide is present a white cloud is produced. A solution of barium hydroxide, Ba(OH)2, can be used instead of one of calcium hydroxide; a similar reaction takes place and barium carbonate, BaCOa, is formed. 206. Uses of Carbon Dioxide. The gas is used in making washing soda, baking soda, white lead, and other carbonates, to aerate water, and in the preparation of sparkling beverages like ginger ale, etc. Portable fire extinguishers which contain a solution of a car- bonate and an acid are much used; a common form is sketched in Fig. 19. The apparatus is nearly filled with a solution of baking soda, and is closed by screwing on a cover to which is attached a frame containing a bottle filled with concentrated sulphuric acid. The neck of the bottle is closed by a loosely fitting stopper of porcelain or glass. This is provided to prevent the acid from absorbing water from the solution; if this occurs the acid overflows and slowly reacts with the carbonate. When the extinguisher is to be used it is inverted, the stopper drops out of the bottle, and the acid mixes with the solution of the carbonate. Carbon dioxide is formed, and rising through the solution, collects above the liquid. The pressure produced forces the water out of the tube at the side. A stream of water containing sodium bicarbonate and carbon dioxide under considerable pressure is available, as a result, for extinguishing fires. The fact that air containing as little as 15 per cent of carbon dioxide will prevent wood from burning is one of the factors in making the apparatus very efficient. 207. An ingenious method of extinguishing burning oils has been recently developed. When water is poured on a burning oil the latter floats and continues to burn. If water containing car- bon dioxide is used, the gas is quickly carried away by the hot products of combustion and does not extinguish the fire. The difficulty was overcome by using two solutions which were mixed FIG. 19. 186 INORGANIC CHEMISTRY FOR COLLEGES and allowed to flow over the burning oil. One solution contained aluminium sulphate and a small amount of glue, and the other sodium bicarbonate. When the solutions are mixed, carbon dioxide is formed, as the aluminium carbonate produced as the result of the double decomposition immediately breaks down into aluminium hydroxide and carbon dioxide. The glue in the water serves to retain the gas in the form of bubbles, which persist for a long time. The foam produced in this way covers the burning oil and extin- guishes it. Fire extinguishers are also provided which contain solid mate- rial to be thrown on a flame; in these, a carbonate is used which at the temperature of the flame decomposes and liberates carbon dioxide. The fact that carbon tetrachloride is used in one type of fire extinguishers has already been noted; in this case when the liquid is forced in a stream into the fire, it changes to a vapor which prevents combustion. 208. Carbon Dioxide in Nature. The fact has been mentioned that plants grow as the result of the transformation of water and the carbon dioxide of the air into the materials of which they are composed; and it has also been pointed out that when plants decay they break down into carbon dioxide, which returns to the air. The changes which take place are not yet fully understood as they occur step by step, but the final products of growth and decay have been carefully studied. The passage of carbon as car- bon dioxide from the air into the materials of plant life, and back to the air again, is sometimes called the cycle of carbon. Through this and the cycle of other elements, such as nitrogen, the living world constantly regenerates itself. The plant is so constituted that when it dies the products of its decomposition serve as food for a growing plant. A forest can regenerate itself forever. In the case of animals, however, the waste materials or the products formed when death occurs, do not serve as food for other animals; they are assimilated by plants, however, and thus return to take part in the cycle of carbon. The relation between plant and animal life is striking; the waste of the animal becomes the food of the plant, which in turn becomes the food of the animal. 209. The energy changes which take place in the cycle of carbon are of great interest. The life-processes of animals are associated CARBON AND ITS OXIDES 187 with oxidation in which energy is set free as heat. The foods we eat pass into carbon dioxide, which is exhaled, and heat is gene- rated in the body. In the life of an animal chemical energy is dis- sipated as heat the store of available energy is thus constantly decreased. On the other hand, when a plant increases in weight through growth, reduction takes place and energy is stored up. When carbon dioxide reacts with water to form cellulose, the sub- stance of which the woody fiber of plants is composed, it under- goes reduction, and oxygen is set free. The energy required is taken up from the sunlight, for the reaction occurs only when this form of energy is available. Plants give off carbon dioxide at night. Complex chemical changes, which have been studied care- fully but are as yet not understood, occur near the surface of the leaves of plants; it is known, however, that the green coloring matter of the leaf, chlorophyll, plays an important part in the transformations as they take place. If a growing plant is placed under water containing carbon dioxide and put in the sunlight, the oxygen generated can be collected and tested. We see then, that sunlight is the source of the energy stored up in plants. When we climb a hill we are using energy from this source, when we run an engine by burning wood or coal, the sun is doing the work. When we use a water-fall to generate electricity we are again obtaining available energy which came originally from the sun; for the heat of the sun evaporated the water from the surface of the earth, and the vapor rose, formed clouds which turned into rain, and finally fell on an elevated part of the earth's surface, and in seeking a lower level the water gave up a part of its potential energy. CARBON MONOXIDE 210. When charcoal smoulders a gas is formed, which is char- acterized by the fact that it is deadly poison; it is called carbon monoxide and has the composition represented by the formula CO. The gases formed as the result of firing explosives contain carbon monoxide; as a consequence, the air in the turrets of battleships may contain the gas after the guns have been repeatedly fired. Special precautions are taken, therefore, to assure the adequate ventilation of the turrets. During the recent war a gas mask was devised that furnishes protection against the gas. Carbon 188 INORGANIC CHEMISTRY FOR COLLEGES monoxide was not used as a war gas on account of its lightness and its relatively low toxicity. Carbon monoxide is an active poison. The gas unites directly with the red coloring matter of the blood, and thus prevents the formation of the compound of the latter with oxygen. It will be re- called that oxygen is carried to all parts of the body by the blood ; if this is prevented death results. The compound formed when carbon monoxide gets into the blood has a bright-red color; it is for this reason that the skin of persons poisoned by the gas has such a pink color after death. Ten cubic centimeters of carbon monoxide per kilogram weight of an animal will cause death; a man of average weight, will be killed by inhaling about 800 c.c. of the gas. Air which contains 1 part of carbon monoxide in 800 will prove fatal in about thirty minutes; as little as 1 part in 2000 causes un- consciousness and convulsions, and, finally, death. The presence of very small amounts of carbon monoxide in the blood can be de- tected by examining by means of a spectroscope (615) the light which passes through it; the blood absorbs a part of the light which passes through unchanged when pure blood is examined. Traces of carbon monoxide have been found in tobacco smoke; it is claimed that the injurious effects of inhaling the gases produced in smoking can be traced to this cause. 211. Preparation of Carbon Monoxide. When carbon dioxide is passed over red-hot zinc or iron it is reduced to carbon monoxide (203) ; it is also formed when carbon is used : CO 2 + C = 2CO This is a reaction of importance and is used in the preparation of producer gas, which is described later (229). Carbon monoxide is formed along with hydrogen when steam is passed over red-hot coal : H 2 O + C = CO + H 2 This reaction is the basis of the preparation of water-gas, which is used for purposes of illumination and as a source of heat and power (228). In the laboratory carbon monoxide is most conveniently pre- pared by treating formic acid with concentrated sulphuric acid; CARBON AND ITS OXIDES 189 the reaction consists in the removal of water from formic acid, the sulphuric acid serving as the dehydrating agent : H 2 CO 2 = H 2 + CO The preparation is carried out by allowing formic acid to fall, drop by drop, from a funnel provided with a stop-cock, into con- centrated sulphuric acid; the latter is contained in a flask con- nected with a delivery tube arranged to collect a gas over water. When oxalic acid is heated with sulphuric acid both carbon monoxide and carbon dioxide are formed: H 2 C 2 O 4 = H 2 O + CO + CO 2 Carbon monoxide is separated by passing the gas through a sol'i- tion of sodium hydroxide; the carbon dioxide is converted into sodium carbonate, which remains in the solution, and the carbon monoxide passes on unchanged. 212. Physical Properties of Carbon Monoxide. Carbon mon- oxide is a colorless gas, with practically no odor or taste. It is slightly lighter than air; 1 liter at and 760 mm. pressure weighs 1.250 grams. Water dissolves at 20 about 2 per cent of its volume of the gas. It has been condensed to a liquid which boils at 190 and solidifies at -203. 213. Chemical Properties of Carbon Monoxide. Carbon monoxide is not active at ordinary temperatures. It burns with a blue flame when ignited : 2CO + O 2 = 2CO 2 The blue flames observed over the surface of coal that is burning slowly under a weak draught are produced as the result of the combustion of carbon monoxide. Carbon monoxide unites with chlorine in the presence of sunlight to form carbonyl chloride (phos- gene) : CO + C1 2 = COC1 2 Carbonyl chloride boils at 8, and is decomposed by water into carbon dioxide and hydrochloric acid; it is used in the preparation of dyes, and was an important war-gas. At high temperatures carbon monoxide is a powerful reducing agent; the gas is formed in the furnaces used to extract metals 190 INORGANIC CHEMISTRY FOR COLLEGES from their ores, and is the chief agent in effecting their reduction. The reactions with copper oxide and ferric oxide are represented by the following equations : CuO + CO = Cu + CO 2 Fe 2 O 3 + SCO = 2Fe + 3CO 2 214. Test for Carbon Monoxide. The fact that carbon monox- ide burns with a blue flame and does not support combustion is used as a test for the gas; this behavior is not characteristic, how- ever, and the conclusion that a gas is carbon monoxide because it acts in this way must be confirmed by other tests. Carbon monoxide dissolves in a solution of cuprous chloride, CuCl, in ammonia. The reagent prepared in this way is used to separate carbon monoxide from other gases, and finds a valuable application in the analysis of flue gases, which contain carbon monoxide. 215. Carbon Bisulphide. When the vapor of sulphur is passed over charcoal or coke at red heat, carbon disulphide, CS 2 , is formed. The reaction is commonly carried out in a furnace in which an electric current is the source of heat. Carbon disulphide is an endothermic compound ; when the liquid is formed the energy change is represented by the following equation: C + S 2 = CS 2 - 19,600 cal. For this reason carbon does not burn in sulphur vapor as it does in oxygen. The ease with which carbon disulphide decomposes, and the low temperature at which it begins to burn (about 200) can be traced to the fact that the molecule contains more chemical energy than the free elements of which it is composed. Carbon disulphide is a colorless liquid which has the specific gravity 1.26 at 20; when pure it has an ethereal odor, but as ordinarily obtained it contains compounds of sulphur which impart to it a very unpleasant smell. Carbon disulphide boils at 46, and vaporizes rapidly at room temperature; the vapor is about 2.7 times as heavy as air and does not, therefore, diffuse rapidly. Owing to this and the fact that the vapor burns when heated with air to about 200, carbon disulphide is a dangerous substance to work with in the neighborhood of flames. As in the case of other inflammable liquids, a mixture of its vapor and air explodes CARBON AND ITS OXIDES 191 when ignited. When carbon disulphide burns, carbon dioxide and sulphur dioxide, a gas with a pungent disagreeable odor, are formed : CS 2 + 3O 2 = CO 2 + 2SO 2 Carbon disulphide is almost insoluble in water, but mixes in all proportions with ether, benzene, alcohol, and many oils. It is an excellent solvent; it dissolves phosphorus, iodine, sulphur, wax, tars, resins, rubber, oils, and fats; on this account it is employed as an extracting agent, but owing to the danger connected with its use it is being replaced when possible by carbon tetrachloride. The vapor of carbon disulphide is poisonous, a fact that leads to its use for exterminating moles, rats, woodchucks, etc.; it is also used as a germicide and insecticide. 216. Carbon Tetrachloride. On account of the fact that carbon tetrachloride is a good solvent for oils and grease and is non- inflammable, it has been recently much used to extract wool, seeds, etc., instead of gasoline or benzine. A preparation to replace gasoline as a cleaning agent in the household is sold under the name of " carbona " which consists of benzine and carbon tetrachloridc in such proportions that the mixture or its vapor will not burn. When carbon tetrachloride is poured on a fire the vapor produced from it prevents combustion. It will be recalled that the presence of but 15 per cent of carbon dioxide in air prevents combustion; it is probable that a small proportion of the vapor o'f carbon tetra- chloride acts in the same way. This fact is utilized in the so-called " pyrene " fire extinguisher, which is of particular value in extin- guishing burning oils. If water is applied, the oil floats on its sur- face and continues to burn, but when carbon tetrachloride is used it mixes with the oil and finally renders it non-inflammable. Carbon tetrachloride is made by passing dry chlorine into car- bon disulphide which contains a little iodine that acts as a cata- lytic agent: CS 2 + 3C1 2 = CCU + S 2 C1 2 The carbon tetrachloride is distilled from the mixture and purified ; the sulphur chloride is recovered and used in vulcanizing rubber. Carbon tetrachloride boils at 77 and has the specific gravity 1.628. 217. Calcium Carbide. When lime, CaO, is heated in an elec- 192 INORGANIC CHEMISTRY FOR COLLEGES trie furnace with carbon, the oxide is reduced and the metal liber- ated unites with the excess of carbon present to form a carbide : CaO + 3C = CaC 2 + CO The preparation is carried out in a furnace of the resistance type (192) ; it is built of fire-brick and lined with carbon, which serves as one electrode. The furnace is filled with coarsely pulverized lime and coke, and rods of carbon are introduced through the top of the furnace to serve as the second electrode. As the calcium car- bide is formed it settles to the bottom of the furnace as a liquid, which is drawn off from time to time. Calcium carbide as ordinarily obtained is a hard, crystalline substance of dark color, but when chemically pure is white. It decomposes in the air as the result of the action of the water- vapor present. With water it decomposes rapidly and calcium hydroxide and acetylene are formed : CaC 2 + 2H 2 O = C 2 H 2 + Ca(OH) 2 Commercial calcium carbide is about 80 per cent pure; it is used to make acetylene and in the manufacture of calcium cyanamide (342). 218. Silicon Carbide. When sand, silicon dioxide, is heated to a high temperature with carbon, a reaction analogous to that between calcium oxide and carbon takes place; in this case silicon carbide, SiC, and carbon monoxide are formed: SiO 2 + 3C = SiC + 2CO The reaction is carried out in a resistance electric furnace like that described in connection with the preparation of graphite (192). Between the ends of the graphite electrodes which enter at the sides of the furnace, is packed a core of granulated coke to serve as the conductor of the current. Around and above this is placed a mixture of sand, powdered coke, and salt, the latter being added to assist in binding the materials together. After the reac- tion has taken place, the silicon carbide is found as a layer of bril- liant black iridescent crystals around the central core of carbon, which contains graphite and unchanged coke. The graphite is produced as the result of the dissociation of the silicon carbide at the hottest part of the furnace. CARBON AND ITS OXIDES 193 Silicon carbide is characterized by its hardness, and its use as a grinding material is based on this fact. It has largely replaced corundum, or emery, for this purpose, and is called in trade car- borundum. It is used to make grindstones, knife sharpeners, etc. EXERCISES 1. (a) Name several ways in which manganese dioxide and charcoal, both in the form of a fine black powder, could be distinguished from each other. (6) Which method could be used most conveniently? (c) How could you distinguish a mixture of equal weights of manganese dioxide and charcoal from pure charcoal and from pure manganese dioxide? (d) How could you separate charcoal from a mixture of it with manganese dioxide? 2. (a) How could you free air from carbon dioxide? (6) How could you determine the percentage of carbon dioxide in a sample of air? 3. How could you obtain pure carbon dioxide from the mixture of nitro- gen and carbon dioxide which results when carbon is burned in air? 4. Would you expect calcium carbonate or copper carbonate to be more readily decomposed by heat? Give a reason for your answer. 5. Calculate the percentage by weight of carbon dioxide in a water solu- tion of the gas at 15 and 760 mm., assuming that no carbonic acid is formed. 6. What weight of calcium carbonate must be dissolved in an acid to furnish 10 liters of carbon dioxide at and 760 mm.? 7. How much lime, CaO, can be prepared from 1 ton of limestone which contains 95 per cent CaCO 3 ? 8. How could you tell the percentage of Na 2 CO 3 in a mixture of Na 2 CO 3 and NaCl? 9. (a) What weight, approximately, of CO 2 must be mixed with the air in a room 3X5X6 meters in order to render the air incapable of support- ing combustion? (6) What weight of washing soda, Na 2 CO 3 , 10H 2 O, would have to be treated with an acid to produce this amount of the gas? 10. What volume of carbon monoxide and hydrogen at and 760 mm. is produced when 1 ton of coke containing 90 per cent carbon reacts with steam? (One pound molecular weight of a gas occupies 359 cu. ft. at and 760 mm.) 11. What volume of CO is obtained when 50 grams of formic acid are de- composed by sulphuric acid? 12. How could you separate the CO 2 and CO formed by decomposing oxalic acid with sulphuric acid, and obtain the two gases in pure condition? 13. Devise a method to determine the amount of each of the following gases in a mixture : N 2 , O 2 , CO, CO 2 . Make use of the solubilities of the gases in different reagents. 14. What volume of air contains just enough oxygen to react with the gases formed when 12 grams of carbon decompose water into carbon monoxide and hydrogen? CHAPTER XVI COAL, COKE, ILLUMINATING GAS, FLAMES 219. Coal ranks first among the natural resources which are utilized by man in producing the necessities and comforts of modern civilization. It is the chief source of heat and power, and is used in extracting metals from their ores. The industrial posi- tion of a nation is determined largely by its supply of coal a fact that has led to wars and has influenced the political history of the world. Petroleum, natural gas, and waterfalls are also a source of power, but they furnish but a very small fraction of the energy required to do the world's work. 220. Coal. Coal is the product of the slow decomposition of vegetable material in the absence of oxygen. The early stages of this decomposition can be seen in marshes where grass, leaves, and boughs of trees are undergoing a change which leads to the formation of peat, carbon dioxide, and methane, CEU. When a stagnant pool in which these changes are taking place is stirred, the bubbles of gas which rise to the surface can be lighted and burn as the result of the presence of methane, which is also called for this reason marsh gas. As this change continues, more and more of the oxygen is removed from the cellulose of which the vegetable mate- rial is made up. Cellulose contains carbon, hydrogen, and oxygen in the proportions represented by the formula CeHioC^, although the actual number of atoms of each element present is unknown, and for this reason its formula is usually written (CeHioOs)*. The removal of oxygen produces products which are known successively as peat, lignite, brown coal, and bituminous, semi-bituminous, and anthracite coal. The final changes which lead to the formation of the various grades of coal have been brought about, in all prob- ability, as the result of geological changes on the earth's surface, which produced heat and pressure in the absence of air. Bitu- 194 COAL, COKE, ILLUMINATING GAS, FLAMES 195 minous, or soft coal, contains carbon, hydrogen, and oxygen approximately in the proportion represented by the formula CaeHboC^, although it is not a compound which has this com- position, but is a mixture of many substances. The comparison of this formula with that of cellulose will indicate that soft coal contains a smaller proportion of hydrogen and oxygen than the substance from which it was formed. Anthracite, or hard coal, contains much less oxygen and hydrogen than soft coal. The material from which coal was formed contained in addi- tion to carbon, hydrogen, and oxygen, compounds in which nitrogen, sulphur, iron, silicon, and other elements were present, and as a result, these elements are found in coal. When coal is heated in the absence of air, more or less volatile material is given off which contains ammonia, NHs, hydrogen sulphide, EkS, cyanogen, C2N2, and other gases, some of which are composed of carbon and hydrogen, and are called hydrocarbons. If the residue, which is called coke, is now heated in the air the carbon is burned and an ash remains, which contains the non-volatile, inorganic material in the coal. The ash of coal is a complex mixture which may con- tain the silicates of calcium, magnesium, and iron, and the oxide and sulphide of iron. 221. The Analysis of Coal. The value of a coal as a source of heat is determined by its so-called calorific value, that is, the heat produced when a definite weight of it burns; and its value for making illuminating gas is determined by the percentage of gases given off when it is heated. In analyzing coal 1 gram of the pow- dered air-dried sample is first heated at 105 to determine the free water present. The presence of water lessens the value of the coal, not only because it will not burn, but also because it absorbs a large amount of heat when it is vaporized. A crucible containing the sample is covered and heated to redness in the Bunsen flame. The loss in weight is a measure of the volatile material given off. The coke that remains is then heated in the presence of air and the carbon burned off, the loss in weight being the result of burning the so-called fixed carbon. The residue is ash. A sample of the coal is burned in a bomb-calorimeter to determine its calorific value. The heat produced is expressed for commercial purposes in British thermal units (B.t.u.) per pound of coal. The relation between this unit and the calorie has been explained in section 153. 196 INORGANIC CHEMISTRY FOR COLLEGES Coals from different localities differ in composition. The figures in the following table are the average percentages obtained from analyses of a large number of samples of the fuels listed. ANALYSES OF FUELS Vola- Water. tile Carbon. Ash. B.t.u. per pound. Matter. Wood 20 49 30 1 7,000 to 9,000 Peat 20 51.6 25 3.2 7,000 to 9,000 Bituminous coal 0.9 27.4 64.1 7.6 11,000 to 14,000 Semi-bituminous coal . . . 0.5 16.7 77.3 5.5 13,000 to 15,000 Anthracite coal 2 4.3 86.5 7.2 12,000 to 14,500 Coke . . 2.5 1.3 86.3 12.4 12,000 to 14,000 222. The Burning of Coal. The products formed when coal burns are determined by the proportion of air used ; if an excess of oxygen is furnished, the carbon is converted into the dioxide, the hydrogen into water, the sulphur into sulphur dioxide, and the nitrogen escapes in the elementary condition. If there is not enough oxygen to form these compounds the carbon is converted in part into carbon monoxide, and some of it may escape as smoke; and the sulphur may pass off with the products of combustion as hydrogen sulphide. In burning coal to produce heat for indus- trial purposes attention is paid to the manner in which the coal is added to the furnace the stoking and to the amount of air admitted to the fire. If coal is thrown on the fire so that it forms a thick layer, a part of it is decomposed by the heat, and the volatile matter produced is driven up the chimney along with small par- ticles of solid, by the hot products of combustion. There is, as a result, a loss of combustible matter and smoke is produced. If not enough air is admitted to the fire, a part of the carbon is converted into carbon monoxide, which escapes unburned. This results in the loss of much heat, since the heat of combustion of carbon to the dioxide is 97,000 calories and to the monoxide only 29,000 calories. If more air is admitted to the fire than is required to burn the coal, heat is lost, because the excess air carries away heat when it passes up the chimney with the products of COAL, COKE, ILLUMINATING GAS, FLAMES 197 combustion. The heat lost in the waste gases may amount to as much as 40 per cent of the calorific value of the coal under poor control of the fire. When large quantities of coal are used it is customary to determine the efficiency with which it is burned by analyzing the flue gases; if they contain carbon monoxide or methane, combus- tion has been incomplete and more air should be used ; if they con- tain too large a percentage of oxygen, the air must be decreased. It is impossible to run a furnace so that the theoretical amount of air is used; it has been found in practice that about twice this amount gives the best results. 223. For domestic purposes anthracite coal is preferred because it can be handled with less skill and attention. It is hard and does not crumble, it yields less volatile matter when heated, and does not produce much smoke; and since it does not melt at the tempera- ture of the furnace clinkers do not form. A large supply of coal can be placed on the fire without smothering it. The rate at which coal burns is determined by the amount of air admitted to it, and this can be regulated by the dampers with which the fur- nace is supplied; one is placed below the grate (A), one over the fire in the door where the furnace is fed (B), and one in the pipe that carries off the products of combustion (C). When a fire is started a strong draft is desired and A and C are left open and B closed; by this arrangement the full current of air passes through the coal. If it is desired to check the fire, A is placed so that but a small amount of air can enter the furnace, and the draft produced by the hot products of combustion rising in the chimney is reduced by partially closing C If an insufficient quantity of air to burn the coal is admitted to the furnace in this way, and B is closed, carbon monoxide, methane, and hydrogen sulphide are formed. The air burns the coal resting on the grate to carbon dioxide, and all the oxygen is consumed. The heat generated raises the next layer of coal to incandescence, and when the carbon dioxide passes over it carbon monoxide is formed C + CC>2 = 2CO. At the high temperature volatile products distill from the coal and since no oxygen is present they escape unburned. If the gases leak through a crack in the dome of the furnace, or in any other way get into the current of hot air which passes around the fire- box and is led through pipes to the rooms heated by the furnace, 198 INORGANIC CHEMISTRY FOR COLLEGES " coal-gas " is introduced into the house. This soon becomes evi- dent as the result of the action of the hydrogen sulphide in the gas on any articles made of silver that are exposed to it; they tarnish as the result of the formation of silver sulphide. If the furnace is used to heat water or steam and this is circulated through the house, coal-gas cannot, of course, be introduced in this way. But even if this is the case and no annoyance is caused by burning the coal incompletely, the method is uneconomical, because when carbon is burned to carbon monoxide only about one-third of the heat which can be produced from the coal is obtained. When coal is burning in an insufficient supply of air and the door over the fire is opened, there is usually a slight puff which is caused by the explosion of the mixture of carbon monoxide and air produced, and the gas then burns over the coal with a blue flame. The evident remedy for this condition is to admit a small amount of air through the upper door of the furnace; this checks the draught through the fire, and thus decreases the rate at which the coal burns, and it furnishes the required oxygen to burn the carbon monoxide and other combustible gases produced. 224. Coke. When coal is heated in the absence of air, the vola- tile materials are driven off and the residue obtained, consisting of carbon and ash, is called coke. Large quantities of coke are used in extracting iron and other metals from their ores. A small supply is obtained as a by-product in making illuminating gas, but the large amount required in the metallurgical industries is pre- pared directly for this purpose. Coke is made either in so-called bee-hive ovens (Fig. 20) or in by-product ovens. The former are dome-shaped ovens of brick, about 10 feet in diameter and 6 feet high, provided with a circular opening in the top through which the coal is introduced and the gaseous products formed in the coking are discharged. There is also an opening on one side at the base of the oven, which serves as an entrance for the air and through which the coke is finally withdrawn. The ovens are built side by side and, in certain districts where large quantities of iron are produced, the rows of ovens are a quarter of a mile in length. An oven is charged with enough coal to make a layer 2 feet thick and the air-inlet so adjusted that just enough air is admitted to keep the coal red hot. The heat that passes through the walls from the two ovens on either side, in which coke is being made, soon starts COAL, COKE, ILLUMINATING GAS, FLAMES 199 the combustion of the coal. By charging alternate ovens at the proper intervals the process becomes continuous. It takes about two days to complete the change of the coal to coke. The weight of the latter obtained is about 60 per cent of the weight of the coal. The process is very wasteful because all the gaseous products produced escape into the air and are burned. FIG. 20. 225. Bee-hive ovens are now being replaced, in part, by by- product ovens, which are so constructed that the volatile material formed in coking is saved. When this material is cooled it yields combustible gases, and a mixture of liquids and solids from which benzene, CeHe, toluene, CyHg, and many other valuable substances are obtained. These compounds are the materials from which explosives, dyes, pharmaceutical chemicals, and other important products are made. As the result of the recent war the produc- tion of by-product coke was markedly increased because the by-products were needed in the manufacture of explosives and other organic compounds, such as dyes, which had been pre- viously imported from Germany. About 70 per cent of the coal used is recovered as coke when ovens of this type are used. About 60 per cent of the gas produced is used to heat the coal and the rest is utilized as a source of power. The valuable liquids and solids are condensed, separated, and purified. 226. A by-product oven is made up of a series of rectangular chambers about 33 feet long, 7.5 feet high, and 20 inches wide, 200 INORGANIC CHEMISTRY FOR COLLEGES which are set side by side with spaces between the ovens through which the hot products of the combustion of burning gas pass. The volatile products escape through pipes placed at the top of each retort. The latter has a door at either end. When the coal has been coked these are opened, and the coke is pushed from the furnace by a mechanical device. The products of a by-product oven are about 70 per cent coke, 14 to 16 per cent gas (about 9000 cu. ft. per ton), 4 to 6 per cent coal tar, which contains benzene, etc., and 0.25 to 0.30 per cent ammonia, which is equivalent to about 20 pounds of ammonium sulphate per ton. 227. Illuminating Gas. Coal-gas is made by heating bitumi- nous coal in retorts made of fire-clay about 8 feet long, 18 inches wide, and 15 inches high. Each unit consists of six or eight Hydraulic Main fo-faru Scrubber Condenser i Gas Holder. fteforf- FIG. 21. retorts set together in what is called a " bench," and each is heated either by burning coke on a grate, or by gas. The gas produced is led away through a pipe connected with each retort ; it is first allowed to cool to separate out the tar, and is then passed through " scrubbers " containing water, which dissolves out the ammonia. It is then brought into contact with slaked lime or a hydrated oxide of iron, which removes sulphur compounds from the gas, and is finally stored in a gasometer. (Fig. 21.) The average yields of the products obtained from 1 ton of coal are 10,000 cu. ft. of gas of 16 candle power, 1400 pounds of coke, 120 pounds of tar, and 5 pounds of ammonia. The composition of the gas is determined by the temperature to which the coal is heated, 1000 to 1300. At the lower tem- perature less gas is obtained, but it has a higher candle power. At the higher temperature the gases which produce light when COAL, COKE, ILLUMINATING GAS, FLAMES 201 they burn are in part converted into others which give little light. The former, called illuminants, consist principally of ethylene, C2H4; acetylene, C2H 2 ; and benzene, CeHe; and the latter are hydrogen and methane, CH4. Coal-gas contains ordinarily about 4 per cent illuminants, 49 per cent hydrogen, 35 per cent methane, 7 per cent carbon monoxide, 1 per cent carbon dioxide, and 4 per cent nitrogen. It has a calorific value of about 600 B.t.u. per cubic foot and the specific gravity 0.43 compared with air. When a gas is required for balloons it is prepared by heating coal at a very high temperature so that the chief constituents of it are hydrogen and methane. 228. Water-gas. When steam is passed over highly heated carbon a reaction takes place which is represented by the following equation : C + H 2 = CO + H 2 - 27,100 cal. This reaction is utilized in making water-gas. Since hydrogen burns without light and carbon monoxide with a blue flame, the gas is " enriched " by mixing with it other gases obtained by heat- ing oils to a high temperature. A number of types of apparatus are in use for carrying out these reactions. In one of these anthra- cite coal or coke is brought to incandescence in a generator by means of a blast of air. (Fig. 22.) The hot gases produced are led through the carburetter, which is built of fire-brick and is filled with a " checker-work " of the same material so piled that the bricks furnish channels for the gas to pass through and present a large surface, which is heated white hot. The gases next pass through the superheater, similarly constructed, into which enough air is admitted to complete the burning of the carbon monoxide formed in the generator. The checker-work here is also heated to incan- descence and the products of combustion allowed to escape from the top. When this condition has been reached the air-blast is cut off from the generator and superheated steam is blown through it. The mixture of carbon monoxide and hydrogen produced passes through the carburetter, to which is admitted a slow stream of naphtha or other oil. This is vaporized and decomposed and passes along with the gas into the superheater where the hydro- carbons are converted completely into gases which do not liquefy 202 INORGANIC CHEMISTRY FOR COLLEGES when the gas cools. The enriched or carburetted gas which issues from the superheater is scrubbed and stored. Since the water-gas reaction, C + H^O = CO -f- H2, absorbs heat, the temperature of the coke in the generator soon falls below that at which the reaction takes place, which is about 1000. The steam is then cut off, air admitted and the process repeated. Air is usually blown for about eight minutes and steam for six minutes, the gas being collected only when the steam is used. One ton of anthracite coal yields about 44,000 cubic feet of carburetted gas, which contains approximately 40 per cent of hy- drogen, 17 per cent of methane, 29 per cent of carbon monoxide, Coke 'Smokestack Generator Carbureter Superheater FIG. 22. 9 per cent of illuminants, 4 per cent of nitrogen, and 1 per cent of carbon dioxide. The gas supplied in certain cities is a mixture of coal-gas and water-gas. 229. Producer-Gas. On account of the fact that gas can be used conveniently as a source of heat, that it produces no ashes, and is very efficient as a source of power when used in an explosive engine, the use of gas on the large scale for industrial purposes has greatly increased in recent years. Since only the calorific value of the gas is of importance for these uses, the process by which it is made is simpler than that employed in the manufacture of illumi- nating gas. In the simplest kind of producer to make this kind of gas sufficient air is blown over glowing carbon to convert it into carbon monoxide. The coal used is supported on a bed of ashes COAL, COKE, ILLUMINATING GAS, FLAMES 203 resting on a grate. When the air first comes in contact with the burning coal, the following reaction takes place: C + O 2 = CO 2 + 97,000 cal. The heat generated is carried along by the nitrogen present in the air and by the carbon dioxide and raises the coal to incandescence. This then reacts with the carbon dioxide as follows: CO 2 + C = 2CO - 39,000 cal. By adding these two equations we arrive at the following, which expresses the final reaction : 2C + O 2 = 2CO + 58,000 cal. A large amount of heat is generated and the reaction can be carried out continuously. The carbon monoxide produced can be used as a. source of power because when it burns heat is given off: 2CO + O 2 = 2CO 2 + 136,000 cal. Since 1 volume of oxygen furnishes 2 volumes of carbon monoxide 2C + O 2 = 2CO and air contains 4 volumes of nitrogen to 1 of oxygen, the relation between the volume of CO and N 2 in the gas should be 1 to 2. That this theoretical relation can be closely reached in practice is seen from the following results of the analysis of a producer-gas made as outlined above: CO, 27.0; N 2 , 55.3; C 2 H 4 , 2.5; C 2 H 2 , 0.4; H 2 , 12.0; O 2 , 0.3; and CO 2 , 2.5 per cent. The hydrocarbons and hydrogen are produced from the volatile matter contained in the coal from which the gas was prepared. Producer-gas made in the way described above is called air-gas. When large quantities of gas are required a different type of pro- ducer is used and air and steam are blown simultaneously over coal. The heat required to bring about the water-gas reaction is furnished by that produced by burning the carbon to carbon monoxide. The " semi-water-gas " made in this way has a higher calorific value than the air-gas because it contains a relatively high percentage of hydrogen and much less nitrogen. An analysis of such a gas gave the following results: CO, 27.0; N, 29.0; CH 4 , 2.0; H 2 , 34; CO 2 , 8.0. Producer-gases of both types have a low calorific value on account of the large percentage of nitrogen which they 204 INORGANIC CHEMISTRY FOR COLLEGES contain, but owing to their ease of production they are cheaper per heat-unit than any other form of gas. 230. The Burning of Gas. The chemical changes that occur in a gas flame can be studied conveniently with a Bunsen burner (Fig. 23) . This consists of a tube to which gas is admitted at one end and which is open at the other, where the issuing gas burns. Above the inlet for the gas there are two holes to admit air, which can be opened or closed. When the air supply is closed, ordinary illuminating gas burns with a luminous flame. The oxygen of the air can come into contact with the issuing gas only along its surface, -1610' FIG. 23. FIG. 24. and as combustion takes place the gas within the cone produced is heated to a high temperature. The illuminants present in the gas, principally ethylene and benzene, are decomposed under these con- ditions into hydrogen and carbon, which becomes heated to such a temperature that it gives off light. That the flame contains carbon can be shown by placing a cold object in it; carbon is deposited in the form of soot. When the holes at the base of the burner are opened, air is drawn in by the rising gases and mixes with them. The flame now consists of three cones, which are indicated in the figure. The lower cone, which is non-luminous, consists of air and unburned gas, and as the mixture rises it is heated by the burning gas which sur- COAL, COKE, ILLUMINATING GAS, FLAMES 205 rounds it, until finally ignition takes place. The boundary between the mixture of hot gases and air and the region where combustion begins, is the surface between the non-luminous cone and second cone, which is blue. In the latter, partial combustion of the gases is taking place. This has been determined by withdrawing the gases from the blue cone and analyzing them. They were found to consist largely of carbon monoxide and hydrogen even when the gas burned was pure methane, CEU: 2CH4 + C>2 = 2CO + 4H2. The outer zone is practically non-luminous, because carbon mon- oxide burns with a light-blue flame and hydrogen burns without the evolution of light. The fact that the first cone contains relatively cool gas can be shown by a simple experiment. A match is pierced near its head by a pin and then suspended in a Bunsen burner; the pin, resting horizontally across the top of the burner, supports the match, which hangs within the tube. The gas is now lighted, and, if there are no currents of air to blow the flame, the match is not ignited for some time. The temperatures reached in different parts of the Bunsen flame are indicated in Fig. 23. When hydrogen burns in pure oxygen the temperature of the flame is about 2500, whereas when it burns in air the temperature reached is much lower. If equal weights of the gas are burned under these two conditions, the same amount of heat is set free in each case; when oxygen is used, the heat generated raises the temperature of the water-vapor pro- duced up to the higher temperature; when air is used, the large amount of nitrogen present is also heated and, as a consequence, the temperature does not rise so high. It is evident that the highest possible temperature attainable when hydrogen or any other gas is burned with air would be reached when the latter contained just enough oxygen to burn the gas. It is impossible to use such a mixture in a Bunsen burner, for the gas would not burn quietly at the mouth of the burner; the mixture would ignite in the tube and the gas would burn where it enters the burner below the air inlets. If too much air is admitted, the flame " strikes back " because the additional amount of oxygen produces a mixture in the tube which is explosive and which is ignited by the gas burning at the mouth of the burner. It is for this reason that only a part of the oxygen required to burn the gas 206 INORGANIC CHEMISTRY FOR COLLEGES is admitted at the base of the burner and the rest is obtained from the air in the neighborhood of the flame. This difficulty has been overcome in two ingenious ways. In the blast-lamp the gas is burned from a tube inside of which 'is placed a second tube through which air is forced (Fig. 25). The gas and air mix at the ends of the tubes and there is no opportunity for the gas to strike back. In this way the amount of air necessary to burn the gas completely can be mixed with it, and a higher temperature can be attained. FIG. 25. 231. A comparatively recent invention which produces the same result in a simpler way is the so-called Meker burner (Fig. 24) . This is supplied with a heavy grid of metal at the mouth of the burner through which the gases pass. When properly adjusted the flame cannot strike back, because the gases are kept below their kindling-point by being in close contact with the metal of the grid, which is comparatively cold. The arrangement is a novel applica- tion of the principle discovered by Davy and applied by him to the invention of the miner's safety lamp. A much larger supply of air can be used with the Meker burner than with the Bunsen burner and, as a consequence, the amount of oxygen mixed with the gas is nearer that required for complete combustion of the gas and the temperature of the flame is higher. 232. Surface Combustion. The rate at which an explosion travels through a mixture of air and an inflammable gas is deter- mined by the proportions of the two present, the rate being the highest when there is just enough oxygen to burn the gas com- COAL, COKE, ILLUMINATING GAS, FLAMES 207 pletely. It is possible to have certain mixtures of gases move so rapidly that their rate of motion is greater than that of the propa- gation of the explosion through them, and if they are ignited at a distance from the orifice from which they are issuing, they will burn and the flame will not strike back. This principle is used in sur- face combustion. The body to be heated is surrounded by some refractory material in granular form upon the surfaces of which the gases burn. In this way combustion takes place where the heat is desired, a higher proportion of air can be used than is possible when an ordinary burner is employed, and a higher temperature can be reached. This method of combustion has been applied to boilers for generating steam, by filling the tubes with granular refractory material upon the surface of which gas is burned. 233. Long-flame Combustion. When an inert gas like nitro- gen or carbon dioxide is mixed with burning gas the temperature of the flame is lowered, as we have seen. The size of the flame is increased at the same time, for owing to the dilution of the gases combustion takes place more slowly, and the region in which the burning is taking place is larger. This interesting scientific fact has been utilized only recently; its application to industrial uses has resulted in marked improvements in the carrying out of cer- tain manufacturing operations. One example is of particular interest. Lime is made by heating limestone CaCOs = CaO + CO2. If this is accomplished by means of burning coal the material near the fire is heated too hot and that at some distance from it is not completely decomposed. This result was overcome in the old-style kilns by mixing fuel with the limestone, but the lime produced contained the ashes formed and was, therefore, impure. In the Eldred kiln a part of the carbon dioxide formed from the limestone is mixed with the air furnished the coal on the grate; carbon monoxide is formed and burns when mixed with the dioxide in a large flame that rises through the kiln and pro- duces the required temperature over a large area. So-called " long-flame " combustion is a marked improvement in the use of gas as a fuel when an even distribution of heat is required. EXERCISES 1. Two samples of coal were analyzed in the way outlined in section 221. The weights given below were obtained. Calculate in each case the per- 208 INORGANIC CHEMISTRY FOR COLLEGES centage of water, volatile matter, fixed carbon, and ash, and state to which class each coal belonged. In the case of one sample the following results were obtained: weight of crucible which contained the coal, 20.000 grams; weight of crucible plus coal, 21.000 grams; weight after heating at 100, 20.986 grams; weight after driving off volatile matter, 20.728 grams; weight after burning fixed carbon, 20.056 grams. The weights obtained with the second sample were, respectively 18.000 grams; 19.000 grams; 18.988 grams; 18.950 grams; 18.069 grams. 2. Why is it that a ton of bituminous coal gives about as much heat as a ton of anthracite coal or coke, although it contains a much smaller per- centage of fixed carbon? 3. (a) Calculate the average weight of the molecules in a mixture con- taining oxygen and nitrogen in the proportion of 1 volume of the former to 4 volumes of the latter. (6) Calculate the average weight of the molecules in coal-gas assuming the average molecular weights of the illuminants to be 40 and using the composition of coal-gas given in section 227. (c) The result obtained in a is approximately the average weight of the molecules of air. From this result and that obtained in 6 calculate the specific gravity of coal- gas compared with air and compare the result with the figures given in the text, (d) Using the composition of water-gas given in section 228 calculate in a way similar to that just used the specific gravity of water gas compared with air. 4. When air-gas is made in a producer how much energy is lost as heat and how much rendered available in the gas produced? Why is it advantage- ous to use a mixture of air and steam in the producer? In this case how much heat is lost and how much rendered available in the producer gas? 5. (a) If a balloon having the weight and size described in problem 12 at the end of Chapter V were filled with coal-gas, what weight would it just lift from the ground? (6) What would be the result if a gas composed of equal volumes of methane and hydrogen were used? 6. (a) Calculate the number of calories in 1 B.t.u. (6) Calculate the factor required to change B.t.u. per pound into calories per gram. 7. Calculate the cost of 1,000,000 B.t.u. produced by burning (a) a sample of anthracite coal costing $12 per ton and furnishing 12,000 B.t.u. per pound and (6.) one of bituminous coal, costing $8 per ton and furnishing 14,000 B.t.u. per pound, CHAPTER XVII ACIDS, BASES, SALTS. SOLUTIONS 234. As the chemical behavior of the substances described in the preceding chapters has been discussed, acids, bases, and salts have been incidentally mentioned. It is advisable at this point to bring together the isolated facts stated here and there, and to discuss more fully in a general way these three classes of com- pounds, as the study of their composition, properties, reac- tions, and uses constitutes the larger part of inorganic chemistry. The attempt to interpret the behavior of these substances when dissolved in water has led to one of the most fruitful theories of modern chemistry a theory which has correlated many facts of prime importance and has given us a much broader and deeper knowledge of chemical phenomena. 235. Metallic Elements. A more or less definite knowledge of what is meant by the word metal has already been gained. We associate the name with substances which possess certain physical properties because these appeal to our senses; metals, when in compact form, have a surface luster which is so characteristic that it is defined by the word metallic; they are, in most cases, hard, ductile, and malleable, and are good conductors of heat and elec- tricity. The chemical properties of metals are also characteristic. They form oxides and hydroxides which dissolve in acids; as the result of the reaction salts and water are formed. The hydroxides of the metals are called bases, and for this reason metals are often called base-forming elements. The more active metals react with acids and hydrogen is set free: Zn + 2HC1 = ZnCl 2 + H 2 Fe + H 2 SO 4 = FeS0 4 + H 2 Both the physical and chemical properties of metals will be dis- cussed more fully later; the facts stated above are sufficient for the understanding of what is to be immediately presented. 209 210 INORGANIC CHEMISTRY FOR COLLEGES 236. Non-metallic Elements. The elements of this class do not possess the properties characteristic of metals. They form oxides which react with water to produce acids, and are, accordingly, called acid-forming elements. Their hydroxides the compounds with hydrogen and oxygen are acids. Chlorine and sulphur are typical non-metallic elements. 237. Bases. The hydroxides of the metallic elements are bases; those which dissolve in water, such as sodium hydroxide and calcium hydroxide, are called alkalies. Sodium hydroxide and potassium hydroxide are caustic alkalies, and they are so called because when left in contact with the skin they " burn." Solutions of bases in water affect the color of many substances, and those which change readily are called indicators. The indi- cators commonly used are litmus, which is changed from red to blue by bases, methyl-orange which changes from red to yellow, and phenolphthalein, a substance which is converted into a red salt by bases. Solutions of bases react with acids and form salts, and as the result the properties of the base disappear. The following are equations for typical reactions : NaOH + HC1 = NaCl + H 2 O Ca(OH) 2 + H 2 SO 4 = CaSO 4 + 2H 2 O 238. Acids. All acids contain at least one hydrogen atom which can be replaced by the more active metals. They have a sour taste, affect indicators, and react with metallic oxides and hydrox- ides to form salts. Acids are classed as monobasic, dibasic, and tribasic, according to the number of replaceable hydrogen atoms they contain; examples are, respectively, HC1, H 2 COs, and HaPC^. Acetic acid, C 2 H4O 2 , contains four hydrogen atoms, but it is monobasic because of these one only can be replaced by metallic atoms; for this reason the formula is sometimes written H C 2 HsO2 or H(C 2 HsO 2 ). Acids may be considered as made up of replace- able hydrogen atoms and acid radicals, which are groups of atoms that remain in combination when the acids react with other sub- stances. The radicals of carbonic acid, H 2 COs, phosphoric acid, H3PO4, and acetic acid, H-C 2 H3O 2 , are respectively, COs, PO4, and C 2 H 3 O 2 . ACIDS, BASES, SALTS. SOLUTIONS 211 239. Salts. When an acid and a base interact to form a salt the reaction is said to be one of neutralization; the resulting sub- stance possesses neither the characteristic properties of the acid nor those of the base. When a solution of an acid is neutralized by one of a base, it is possible to tell when the solution is neutral if an indicator is present; at first the color produced by the acid is seen; as the base is added a point is finally reached when the addi- tion of a single drop causes a per- manent change in color. The solu- tion is neutral when a drop of the acid will produce one color of the 4 are acid salts. A number of methods of naming acid salts are in use. Baking soda, for example, NaHCOs, may be called acid sodium carbonate, sodium bicarbonate, or sodium hydrogen carbonate. When an acid is tribasic it may form three salts with one base. Phosphoric acid forms salts of the composition NaH2PO4, Na2HPO4, and NasPCU; the first is primary sodium phosphate or monosodium phosphate, the second is secondary sodium phosphate or disodium phosphate, and the third is tertiary sodium phosphate or trisodium phosphate. Basic salts usually possess compositions which are more or less complicated. A neutral salt contains a metal and an acid radical; an acid salt contains hydrogen in addition to these; a basic salt contains a metal, an acid radical, and the hydroxyl group, the group characteristic of bases. The following formulas represent the composition of some basic salts: Fe(OH)SC>4, Pb 3 (OH)2(CO 3 )2, and Cu 2 (OH) 2 CO 3 . 242. Solutions. It has already been indicated that acids do not react with metals in the absence of water. The characteristic ACIDS, BASES, SALTS. SOLUTIONS 213 properties of acids due to the hydrogen atom they contain, dis- appear in the absence of a solvent. Many reactions in which bases and salts take part are dependent on the presence of some liquid. Before attempting any explanation of these important facts it will be well to describe some striking experiments to illustrate the behavior of the solutions of several typical substances when sub- mitted to the action of an electric current. A number of vessels are provided which contain, respectively, water, alcohol, kerosene, pure acetic acid, and solutions in water of sugar, hydrochloric acid, sulphuric acid, sodium hydroxide, calcium hydroxide, sodium FIG. 27. chloride, and copper sulphate. An apparatus is provided to determine whether the substances contained in the several beakers will allow electricity to pass through them (Fig. 27) . This consists of an electric lamp (a), one terminal of which is connected with a source of electricity (e) and the other with a piece of platinum foil (6) to serve as an electrode. The latter is attached to a non-conducting stand to which is joined a second piece of platinum (c) to serve as the other electrode, and this is in contact with a wire which returns to the source of electricity. When the apparatus is connected up no current flows on ac- count of the break in the circuit at the electrodes. If a vessel (d) containing a solution which conducts electricity is placed so that the electrodes dip into the solution, a current can pass and the lamp (a) will glow. The substances enumerated above are tested one after the other, the electrodes being wiped dry after each test. It will be found that there is no evidence that either water, alcohol, 214 INORGANIC CHEMISTRY FOR COLLEGES kerosene, or pure acetic acid conducts the current; the lamp does not glow. Neither does the presence of sugar in the water produce any effect; but all the other solutions allow the current to pass and the lamp glows brightly. The substances selected for the ex- periment are typical and represent pure compounds and solutions in water. The conductivity of thousands of substances has been studied and the results are in accord with those obtained in the experiment just described. Solutions of acids, bases, and salts conduct the electric current freely, whereas pure compounds or solutions of other substances such as sugar do not conduct at all or to a very slight degree. Similar results are obtained if certain solvents other than water are used. Although solid acids, bases, and salts do not conduct the electric current, some of them be- come conductors when they are in the molten condition. Another instructive experiment can be performed to emphasize the effect of the presence of water. It was shown that the lamp did not glow when the electrodes were placed in water or in pure acetic acid, which is a colorless liquid. The vessel containing the acid is placed so the electrodes dip into the acid; no current flows. Water is now poured into it slowly; soon the lamp begins to glow very faintly ; as water is added it grows brighter and brighter and, finally, almost as much light is given off as when the current passed through the solution of sulphuric acid. These results are of great interest and need an explanation. Water does not conduct the current freely, nor does acetic acid; but the mixture does. The solution must contain something which is not present in either, before they are mixed. What happens when an acid, salt, or base dissolves in water? A few additional facts must be cited before an answer is given. When an electric current passes through a solution of an acid, base, or salt a profound change takes place; we have seen, for example, that hydrochloric acid is decomposed into hydrogen and chlorine. Electrolysis takes place, and we say that the current passes as the result of electrolytic conduction. A current flows through a metal without producing any chemical change; in this case it passes as the result of what is called metallic conduction. If electrolytic conduction in a solution of a salt is studied, it will be found that all of the metal can be set free at one pole and all the acid radical at the other. Take the case of copper chloride. At ACIDS, BASES, SALTS. SOLUTIONS 215 first the salt is uniformly distributed throughout the solution; at the end of the electrolysis all the copper has been deposited on one pole and all the chlorine set free at the other. It is evident that the metallic part of the salt is attracted by one electrode which is charged negatively, and the chlorine by the other which is charged positively. It is the same with other salts; the metals and the acid radicals appear to be independent and to behave dif- ferently when in solution; and they are attracted by the elec- trodes which are charged with electricity. 243. The Electrolytic Dissociation Theory. Arrhenius, a Swedish chemist, studied critically in connection with the proper- ties of solutions the facts noted above and many others, and showed that they could be interpreted by means of a simple hy- pothesis. It was necessary to explain the fact that solutions of acids, bases, and salts conduct the electric current, that in these solutions the metals and hydrogen behave differently from the acid radicals and the hydroxyl group, but that they all are attracted by electrodes charged with electricity. The assumption made was that the molecules of the substances which dissolved the solute broke down into smaller particles, called ions, and that these, as a result, became charged with electricity. According to this view, when a molecule of hydrogen chloride is dissolved in water, the latter dissociates the molecule into a hydrogen atom which becomes positively charged with electricity, and a chlorine atom which becomes negatively charged. This is represented in symbols thus: HCI -> H + + cr The reasonableness of this view becomes apparent from the fol- lowing considerations. When any two different substances are brought intimately into contact and are then separated, they each take up a charge of electricity, one becoming positively and the other negatively charged. If a piece of silk is brought into close contact with a rod of sulphur by rubbing, and the two are sep- arated, both are found to be charged with electricity. An experi- ment to illustrate this fact has already been described (7). It is not unreasonable, therefore, that in some similar way the separa- tion of two atoms may produce electricity on them. The water brings about the separation, and as it offers a resistance to the 216 INORGANIC CHEMISTRY FOR COLLEGES tendency of positive and negative electricity to unite, the charged atoms remain apart. In general, the greater the resistance shown by a liquid to this tendency, the more it dissociates sub- stances. The atoms charged with electricity which are formed in this way are called ions. The ion of chlorine is assumed to be an atom of chlorine com- bined with a certain amount of electricity; it differs in energy from a simple chlorine atom and should have different properties from those of the latter. We have seen that the amounts of energy associated with carbon in its two forms, diamond and graphite, are different. We are, accordingly, not surprised to find that a chlorine ion possesses different properties from chlorine gas, and it is entirely reasonable that when electricity is removed from the atom and its energy thereby changed, the element should exhibit the properties with which we are familiar. The hypothe- sis put forward by Arrhenius has served to explain so many appar- ently unrelated facts that it is now one of the most fruitful theories of modern chemistry. It is called the electrolytic dissociation theory, because, as the name implies, it involves the assumption of a dissociation of molecules into particles charged with electricity. Sugar when dissolved does not conduct the electric current; accordingly, it is called a non-electrolyte, and does not dissociate into ions. Acids, bases, and salts are called electrolytes because their solutions in water conduct an electric current. The ions of monobasic acids are hydrogen and the acid radical, thus the ions of HC1 are H + and Cl~, of nitric acid H + and NOs". The ioniza- tion of dibasic and tribasic acids takes place in steps. Sulphuric acid dissociates as follows: +H + + HSO 4 ~ HS0 4 --H + + S0 4 ~ The complete ionization of the acid is represented thus: H 2 SO 4 .->2H + + SO 4 ~ It should be noticed that each hydrogen atom carries one positive charge and the number of charges on any atom or group is the same as its valence. The PO 4 ion from phosphoric acid, HaPC^, is ACIDS, BASES, SALTS. SOLUTIONS 217 represented thus, PO 4 . When any substance undergoes ioniza- tion the number of positive charges always equals the number of negative charges a fact which should be carefully noted when ionic formulas are used in writing equations. Bases produce metallic ions and hydroxyl ions : NaOH - Na+ + OH~, and Ca(OH) 2 -> Ca + + + 20HT The ions of salts are formed from the metallic atoms and the acid radicals : NaCl->Na + + Cl~ Na 2 SO 4 ~* 2Na + + SO*" CaSO 4 -*Ca + + + SO*" It is noted again that the number of charges equals the valence of the atom or group; calcium is bivalent and its ion is accordingly Ca ++ . Metals and hydrogen form positive ions, and acid radicals and the hydroxyl group form negative ions. For this reason metals are sometimes spoken of as positive elements and non- metals as negative elements. From the standpoint of the electrolytic dissociation theory the characteristic properties of an acid are due to the hydrogen ion, for these properties are evident only when the acid is in solu- tion and when ions are formed. According to this view an acid is a substance which yields one or more hydrogen ions when it is dis- solved in water. The characteristic properties of bases are due to the hydroxyl ion, OH~; consequently, a base is described as a substance which yields a hydroxyl ion when dissolved in water. 244. Interpretation of Typical Reactions with the Use of the Theory of Ions. Electrolysis. When hydrochloric acid is dis- solved in water it is assumed that it breaks down into hydrogen ions, which are positively charged atoms, and chlorine ions, which are negatively charged atoms. The terminals of a source of elec- tricity, a storage battery, for example, are charged, one positively and one negatively. When these are connected with the elec- trodes (a) and (6) (Fig. 28) which dip into a solution of hydro- chloric acid, we have in contact with the solution two poles that are charged with electricity. The positive charge on a attracts the 218 INORGANIC CHEMISTRY FOR COLLEGES ^ : ~^ negatively charged chlorine ions indicated by in the diagram. As each ion comes in contact with the electrode it is discharged, for when positive and negative electricity come together, the elec- tricity disappears. As the result, in this case, the chlorine ion loses its charge and changes to an atom from which molecules of chlorine are formed; and the gas as we know it is set free. At the same time positive electricity flows from the battery to take the place of that which disappeared when it united with the negative charge on the chlorine ion. As the result of the continual discharge of the electrode in this way, chlorine gas is formed, and the battery furnishes electricity to the pole; as a conse- quence, a current passes through the wire connecting the battery with the electrode as indicated by the arrow. Similar changes take place at the other pole. The positively charged hydrogen atoms give up their charge to the negative electrode; hydrogen is set free, and the battery keeps the electrode negatively charged. As the electrolysis proceeds the negative ions move to the positive pole, which is called the anode and the positive ions to the negative pole, the cathode, and finally all the ions are discharged and the current ceases to flow. The changes described have been studied quantitatively and generalizations of the greatest importance have been dis- covered; this will be considered at length later. When an electric current is passed through a solution of sulphuric acid, hydrogen and oxygen are obtained (41). In this case a secondary reaction takes place at the anode. The negative ion, SC>4~~, is discharged, but unlike chlorine it does not exist as a chemical compound; it immediately reacts with the water, 2SO 4 + 2H 2 O = 2H 2 SO 4 + O 2 , and oxygen is set free. A similar reaction occurs in most cases when an acid radical which cannot exist hi the free condition is liberated ; it changes into the acid by uniting with hydrogen, and oxygen is set free. The acid is thus FIG, 28. ACIDS, BASES, SALTS. SOLUTIONS 219 continually regenerated. It was stated that in the decomposition of water into hydrogen and oxygen, sulphuric acid acted as a catalytic agent (27); the manner in which it acts is now clear. The hydrogen and oxygen come from the water but the latter first reacts with the discharged SO 4 ion and forms sulphuric acid, which is decomposed by the current. 245. Neutralization. The following equations, in which the ions involved are represented, bring out relationships in neutraliza- tion which would not be apparent without the theory of electro- lytic dissociation. Na + + OH- + H + + OP = Na + + CP + H 2 O K + + OH- + H + + cr = K + + cr + H 2 o C a + + + 20H~ + 2H + + 2C1~ = Ca ++ + 2C1~ + 2H 2 O Ca ++ + 20H- + 2H + + SO 4 ~"~ = Ca ++ + SO 4 ~" + 2H 2 O In each case the reaction consists in the union of' hydrogen ions with hydroxyl ions to form molecules of water, which are undis- sociated. A sodium ion, Na + , is represented on the left-hand side of the first equation; it was formed when sodium hydroxide was dissolved in water. A sodium ion is represented on the right-hand side of this equation. From the viewpoint of ions the sodium has not taken part in the reaction at all; likewise the chlorine ion has been inactive. We do not write NaCl on the right side of the equation to indicate that sodium chloride is formed ; in solution the salt is ionized and should appear as Na + + Cl~. Whether such a view is correct can be tested experimentally. In the reactions represented by the first two equations above the only chemical change is the formation of water from hydrogen and hydroxyl ions; if this is true the heat produced when the quan- tities represented in the equation are brought together should be the same. If 1 gram-molecule of sodium hydroxide (23 + 16 + 1 = 40 grams) is dissolved in water and neutralized with a solution of hydrochloric acid (1 + 35.5 = 36 grams), 13,600 calories are evolved. If a similar experiment is carried out and 1 gram-mole- cule of potassium hydroxide (39 + 16+1 =56 grams) is used, the same amount of heat is set free. When the heat evolved in the reactions represented by the third and fourth equations above is 220 INORGANIC CHEMISTRY FOR COLLEGES determined in the same way, it is found that 2 X 13,600 calories are set free in each case, for 2 gram-molecules of water are formed. The heat of neutralization is a constant and can be represented thus: H + + OH~ = EkO + 13,600 cal. This statement holds true only when solutions containing com- pletely dissociated acids and bases are used. We shall see later that under certain circumstances acids, bases, and salts are only partially broken down into ions. The fact that the heat of neu- tralization is a constant is strong evidence of the correctness of the conception of ions. 246. Acids , bases, and salts have each two distinct sets of prop- erties. We have seen that certain properties of acids can be attributed to the hydrogen ions which they yield when dissolved in water. Each acid also possesses properties which are deter- mined by the acid radical it contains; thus a solution of hydro- chloric acid contains hydrogen ions, H + , and chlorine ions, Cl~~, and the properties of the latter are shown by the acid. Likewise sodium chloride gives a sodium ion and a chlorine ion, when dis- solved in water. If this view is correct its solution should possess the same set of properties which hydrochloric acid possesses due to the fact that it, also, yields a chlorine ion. Or making the conclusion more general, all substances, whatever they may be, should show certain properties in common provided they all yield a chlorine ion when dissolved in water. Such a conclusion is a logical consequence of the theory and the facts are in accord with the theory. We have learned that a test that can be applied to all the chlorides of the metals is to treat their solutions with a solution of silver nitrate; in all cases silver chloride is formed they all show one property in common. Ionic equations for the reactions serve to emphasize this fact : K + + Cr + Ag + + NO 3 ~ = AgCl + K + + N0 3 ~ Ag + + NO 3 ~ = AgCl + Na + + NO The reaction in the two cases represented above consists in the union of the silver and chlorine ions to form silver chloride, which separates as a precipitate; this occurs as the amount of the chlo- ACIDS, BASES, SALTS. SOLUTIONS 221 ride that can be formed ordinarily when solutions are brought together is greater than the amount which will remain dissolved in water (1 liter of water dissolves 0.0013 gram AgCl). Silver chlo- ride is written in the non-ionized condition as AgCl, because it separates as a precipitate; ions exist in solutions only, therefore in writing ionic equations, substances not dissolved are expressed by their molecular formulas. From the above it is seen that the test for chlorides (145) is a test for chlorine ions; it can be made not only with silver nitrate, but any soluble silver compound which gives a silver ion. Like acids, salts and bases when in solution have two distinct sets of properties; all sodium salts, for example, have a set of properties in common upon which a test for sodium can be based. The facts enumerated above find a satisfactory explanation in the theory of ions; they were known before the theory was put for- ward, but no reasonable interpretation of them was given. 247. Reactions of Double Decomposition. The application of the theory of electrolytic dissociation to reactions of neutralization and to the test for chlorides can be broadened to include all other cases of double decomposition. A few equations written with ionic symbols will make this clear: 2Na + + CO 3 ~ + Ca ++ + 2C1" = 2Na^ + 2C1" + CaCO 3 Fe + + + 2Cr + 2Na + + 2OH~ = 2Na + + 2C1~ + Fe(OH) 2 2K+ + 20H~ + 2H + + SO 4 = 2K + + SO 4 " + 2H 2 O 2Na + + 2Cr + 2H++ SO 4 ~~ = 2Na + + SO 4 ~" + 2HC1 In all cases of double decomposition at least one substance escapes from the system undergoing change. In the equations given above CaCOs and Fe(OH)2 are written in the molecular form because they are insoluble substances and precipitate; they are not in solution and, therefore, not in the ionic condition. The formula of hydrochloric acid is written in the molecular form because it escapes as a gas. The formula for water is also molecular because it does not appreciably dissociate. It has been stated (149) that a double decomposition takes place if one of the products is a gas or is insoluble. A third condition can now be added; reactions of this type take place if one of the products is a substance which is un- 222 INORGANIC CHEMISTRY FOR COLLEGES dissociated or slightly dissociated. It is seen from the above that if we wish to make any insoluble acid, base, or salt we can do so by bringing together in solution two substances one of which yields the positive and the other the negative ion of the compound sought. 248. The Reaction between Metals and Acids. The reaction between zinc and sulphuric acid is represented with ionic symbols as follows: Zn + 2H + + S0 4 ~" = Zn^ + H 2 + S0 4 ~ Hydrogen gas, H2, and zinc sulphate, ZnSC>4, are formed; as the salt is dissolved in the water used as a solvent it is represented as dissociated into ions. The only change which takes place is, according to this view,the transfer of the electric charge from the hydrogen to the zinc. When copper is treated with a solution of sulphuric acid a similar change does not take place; this is explained by saying that hydrogen has a greater tendency to be an ion than copper, and does not give up its charge to the metal. With zinc, however, the case is different; it has a greater tendency to be an ion than hydrogen and, accordingly, passes into solution and forces the hydrogen out. All elements tend more or less to pass from the condition of the free element into that of an ion; they differ markedly among themselves in the strength of this tendency, which determines the chemical behavior of the elements in many of their reactions with other substances. 249. The Reaction between Metals and Salts. In the cases cited the tendency of zinc to form ions is greater than that of hydrogen, and that of copper is less than that of hydrogen. We might well ask what would happen if metallic zinc were put into a solution of copper sulphate. The experiment can be readily performed, and the results are what is anticipated ; they are represented by the following equation : Zn + Cu++ + S0 4 -- = Zn++ + Cu + SO 4 The greater tendency of zinc to form ions asserts itself; it passes into solution and metallic copper is precipitated. 250. The Reaction between Hydrogen and Salts. Another case of replacement of this kind is of interest. Platinum, we have learned, is a very inactive metal; it forms ions less readily than ACIDS, BASES, SALTS. SOLUTIONS 223 hydrogen and, consequently, does not replace it when the metal is brought into contact with an acid. We might prophesy what would happen if hydrogen gas were passed through a solution of platinum chloride, in which platinum ions and chlorine ions are present. The greater ionizing tendency of hydrogen comes into play; it becomes the ion and the platinum is forced out of solu- tion. The equation for the reaction is as follows: Pt ++ + 2C1" + H 2 = Pt + 2Cr + 2H + When hydrogen gas is bubbled through a solution of platinous chloride, platinum is precipitated. 251. The Reaction between Non-metals and Salts. One more case of replacement, in which negative ions are involved, will serve to emphasize the principle under discussion. It will be recalled that in one of the tests for free chlorine the gas is treated with a solution of potassium iodide (130); iodine is set free and is recog- nized by its action on starch. According to the conception of ions the reaction takes place as indicated by the following equation: C1 2 + 2K + + 2I~ = 2Cr + 2K + + I 2 When chlorine passes into solution it becomes a negative ion, since it is a non-metallic element. Its tendency to pass into the ionic condition is greater than that of iodine, and when brought into contact with iodine ions the exchange in the electric charge takes place, and iodine is set free in the elementary condition. Elaborate investigations have been made of the replacing power of elements, because a measure of the tendency of an element to pass into the ionic condition is an index not only of its ability to replace other elements, but of its general chemical activity. The elements which form ions readily are the ones we have called active elements, because a large amount of energy is set free when they react with other substances. 252. Electromotive Series of the Metals. There are a number of independent ways of measuring the tendency of a free element to pass into solution as an ion, and they all give approximately the same result. One method is to place samples of a given metal into solutions of salts of other metals and determine if a reaction takes place. For example, if some metallic mercury is placed in a solution of silver nitrate it will be found that the former passes 224 INORGANIC CHEMISTRY FOR COLLEGES into solution and silver separates out ; the tendency of mercury to to form ions is, therefore, greater than that of silver. By a series of experiments it is possible to find out which metals possess this property to a greater or less degree than any given metal. By carrying out experiments in this way with each metal it is pos- sible to make a table in which the elements are arranged according to their relative tendencies to form ions. There is given for reference on this page such a table, in which the arrangement is according to the decreasing values for the com- mon metals. It will be seen that what we have called the active metals are at the top of the list. All the metals above hydrogen set it free from acids, whereas those below hydrogen do not. We shall see as the chemistry of the metals is studied that the position of an element in the series is, in general, an indication of its chemical behavior. This arrangement of the metals is called the elec- tromotive series because it was arrived at as the result of the study of certain electrical properties of solutions which will be discussed later. 253. The Boiling-point of Solutions. Pure water boils at 100 at 760 mm. If a solid is dissolved in it, it will be found that the solution must be heated to a higher temperature before it boils. For example, the boiling-point of water can be raised to 108 by dissolving salt in it. The surface of the solution contains a smaller number of molecules of water than an equal surface of pure water. In order to have the same number of molecules leave the two surfaces in the same time and thus produce the same vapor-pressure (179), it is necessary to raise the temperature of the solution to impart a greater velocity to the molecules of water in the solution. The effect of dissolved substances on the boiling-points of liquids was carefully investigated by Raoult, a French scientist, who discovered facts of the greatest interest and importance. He found that there appeared to be no regularity in the behavior of solutions of acids, bases, and salts, but that the effect on the boiling- point of what we now call non-electrolytes was such that it could be Potassium Sodium Barium Strontium Calcium Magnesium Aluminium Manganese Zinc Chromium Cadmium Iron Cobalt Nickel Tin Lead Hydrogen Copper + + Arsenic Bismuth Antimony Mercury + Silver Palladium Platinum Gold ACIDS, BASES, SALTS. SOLUTIONS 225 expressed by two simple laws. The first of these states that the rise in boiling-point of a solution of a given non-electrolyte is pro- portional to the weight of the dissolved substance provided the same weight of solvent is taken in each case. For example, when 100 grams of glycerine are dissolved in 1 liter of water the boiling- point of the solution is 100.56; if 50 grams are dissolved in the same amount of water the boiling-point is 100.28, the rise being one-half that produced by 100 grams. The second law is as fol- lows: The boiling-points of all solutions of non-electrolytes in the same solvent are the same, provided the amounts dissolved in a given weight of the solvent are in the ratio of the molecular weights of che solutes. The molecular weights of sugar, Ci 2 H 2 2On, glycerine, CaHgOs, and glucose, C 6 Hi2O 6 , are respectively 343, 92, and 180. If 343 grams of sugar are dissolved in 1000 c.c. of water the solution will boil at 100.52; if 92 grams of glycerine or 180 grams of glucose are dissolved in 1000 c.c. of water the boiling- point in either case will also be 100.52. Neither of the above laws holds for electrolytes; for example, a solution which contains 1 gram-molecule of salt, NaCl (58.5 grams), dissolved in 1000 c.c. boils at 100.87. No explanation of " the fact that the behavior of solutions of acids, bases, and salts is not in accord with these laws was furnished until the electrolytic dissociation theory was put forward. We shall see that this theory supplied not only a reasonable interpretation of the observed results, but opened up the way for a detailed study of solutions from a new point of view, which has increased markedly our knowledge of chemistry. If the weight of a molecule of a certain substance is 30 and that of another is 60, it is evident that 30 grams of the first contain as many molecules as 60 grams of the latter. Solutions of non- electrolytes which contain in the same amount of solvent the weights of the dissolved substances proportional to their molecular weights, have the same boiling-point; it follows, therefore, that solutions which contain the same number of molecules in the same volume boil at the same point. With these facts in mind the behavior of solutions of electrolytes can be understood. According to the electrolytic dissociation theory molecules of these substances break down into ions; a solution of sodium chloride, for example, if completely ionized, contains twice as 226 INORGANIC CHEMISTRY FOR COLLEGES many particles in solution as it would contain if it were not ionized : NaCl -> Na + + Cl~ If we assume that the rise in the boiling-point of a solution is deter- mined not by the number of molecules in solution, but by the number of particles, molecules or ions, the behavior of acids, bases, and salts finds a ready explanation. If dilute solutions of sugar and salt are made by dissolving in equal amounts of water a gram- molecule of each, it will be found that the rise in the boiling-point of the salt solution is just twice that of the sugar solution; each molecule of salt breaks down into two ions, whereas the molecule of sugar does not change. As a consequence, one solution contains twice as many particles as the other and the rise in its boiling-point is twice as great. When 1 gram-molecule of salt is dissolved in 1000 c.c. of water the rise in the boiling-point is 0.87; if no dissociation had taken place the rise would have been 0.52, and if it had been complete it would have been 1.04. Evidently but a part of the molecules undergo dissociation at this concentration. The fraction which breaks down into ions can be calculated as follows: Consider 100 molecules and let a equal the number which are converted into ions ; a is, then, the percentage dissociated. The number of un-ionized molecules is expressed by 100 a; and since in the case of sodium chloride two ions are formed from one molecule, 2a represents the number of ions. The total number of particles formed from 100 molecules is then 100 a + 2a, which simplified is 100 + a. Now, since the rise in the boiling-points of solutions is proportional to the number of particles present, we have: 0.52 ? ~ 0.87' ] The salt is, then, dissociated to the extent of 67 per cent. It has been found that the calculation of the degree of dissocia- tion by means of the determination of the rise in the boiling- point of solutions gives results which agree with those arrived at by other methods only when dilute solutions are used. When weights greater than one-tenth of a gram-molecule are dissolved in a ACIDS, BASES, SALTS. SOLUTIONS 227 liter of water deviations are observed, which grow greater as the concentration is increased. 254. The Freezing-point of Solutions. When substances are dissolved in water and other liquids, the freezing-points of the solutions are lower than those of the pure solvents. Laws similar to those stated above in regard to the elevation of the boiling-point have been found to express the influence of dissolved substances on the freezing-points of solutions. The effect on the freezing- point is greater than on the boiling-point. A molecular weight of a non-electrolyte dissolved in 1000 grams of water raises the boiling-point 0.52, but lowers the freezing-point 1.86. The extent of dissociation can be calculated from the freezing-point in the same way as that illustrated above in the case of the boiling-point. When dilute aqueous solutions freeze, pure water usually sepa- rates out in the solid form. In order that the ice may be in equi- librium with the solution, the temperature must be that at which the melting of the ice and its formation from the solution take place at the same rate. This temperature is below 0. 255. A saturated solution of salt freezes at 21, which is 6 below zero Fahrenheit. This means that at this temperature pure ice and a saturated solution of salt are in equilibrium. It will be recalled that the freezing-point is denned as the temperature at which the solid and liquid phases are in equilibrium (187), and that at equilibrium an addition of heat does not cause a rise in tempera- ture; some of the solid changes to liquid. If solid salt is put on ice at 0, some of the former dissolves and a saturated solution of salt is formed; the amount which dissolves may be vanishingly small, but the solution which results can be in equilibrium with ice at 21 only. In order that the temperature may fall, some of the heat present in the ice which has the temperature must dis- appear; the ice melts and in so doing absorbs heat and the tem- perature falls. As the ice melts more salt dissolves in the water, and the melting continues until the temperature is 21. These facts are utilized in making freezing-mixtures for freezing ice-cream and in keeping switches on car tracks free from ice. A mixture of 100 parts of calcium chloride, CaCk, GH^O and 70 parts of snow will lower the temperature to 54. 256. Determination of Extent of lonization from the Conduc- tivity of Solutions. The extent to which an acid, base, or salt 228 INORGANIC CHEMISTRY FOR COLLEGES ionizes can also be determined by a study of the conductivity of solutions. We have learned that the best explanation which has been offered of the way in which a solution conducts electricity is based upon the assumption that ions are present in the solutions (243) . If this is true the larger the number of ions present the more readily will the solution conduct. The extent to which an elec- trolyte breaks down into ions is determined by the amount of water present. The experiment with acetic acid (242) showed that the acid alone did not conduct the current, but its solution in water did conduct, and that the conductivity increased more and more as water was added and the solution was diluted. When a similar experiment is made with hydrochloric acid or salt, it is found that a point is finally reached when the addition of more water has no effect on the conductivity; at this point the molecules are assumed to be completely dissociated into ions. In order to find out the extent to which a salt is dissociated into ions at any con- centration, we first determine the conductivity of its solution, when completely ionized, in an apparatus arranged so that all the ions take part in carrying the current, and then the conductivity of the solution having the special concentration. Suppose in the first case the conductivity is represented by the number 10 and in the next case by 6 then 0.6 of the salt was ionized in the second case. 257. Methods of Expressing Concentration of Solutions. It has been stated that the extent of the ionization of electrolytes is determined by the relation between the amounts of the solvent and the dissolved substance. This relation the concentration can be expressed in a number of ways; one is to state the number of grams dissolved in a liter of the solution; another method, which is particularly valuable when solutions of various substances are to be compared, is to express concentration in gram-molecular-weights per liter. A solution which contains 1 gram-molecular-weight of the dissolved substance in 1 liter of the solution is called a molar solution; if one-tenth of this amount is present the solution is 0.1 molar. When acids are compared another method is often used. The characteristic properties of acids are due to the presence of the replaceable hydrogen atoms they contain. A molecule of hydro- chloric acid, HC1, is not equivalent to one of sulphuric acid, H2SO4, because the latter contains two replaceable hydrogen atoms; 1 ACIDS, BASES, SALTS. SOLUTIONS 229 gram-molecular-weight of sulphuric acid will neutralize twice as much sodium hydroxide as 1 gram-molecular-weight of hydro- chloric acid. For this reason another method is used in expressing the concentration of acids. A normal solution of an acid is defined as one which contains in 1 liter of solution the amount of the acid that furnishes 1 gram of replaceable hydrogen. A normal solution of hydro- chloric acid and all monobasic acids contains 1 gram-molecule in a liter. A normal solution of sulphuric acid and all dibasic acids contains one-half of a gram-molecule in a liter. A normal solution of phosphoric acid, HsPCU, contains one-third of a gram-molecular- weight in a liter, etc. It is evident that a definite volume of a normal solution of any acid is equivalent to the same volume of a normal solution of any other acid. If 25 c.c. of a normal hydro- chloric acid solution is required to react with a certain weight of zinc or to neutralize a definite amount of sodium hydroxide, then 25 c.c. of a normal solution of any other acid will react with these amounts. The concentrations of solutions of bases are expressed in a similar way. A normal solution of a base which has one hydroxyl group, such as NaOH, contains 1 gram-molecular-weight in a liter. A normal solution of a di-acid base, like Ca(OH)2, contains one- half of a gram-molecule in a liter. It is evident that 1 c.c. of a normal solution of any base will neutralize 1 c.c. of a normal solu- tion of any acid. A normal solution of an oxidizing agent contains enough of the compound dissolved in a liter to oxidize 1 gram of hydrogen, and a normal solution of a reducing agent contains in a liter that amount of the substance which equals in reducing power 1 gram of hydrogen. The concentration of any solution can be expressed as a fraction or multiple of a normal solution; we can have, for example, a twice normal, a one-tenth normal, a 0.95 normal and a ^j normal solution of hydrochloric acid, etc.; these concentra- tions are expressed thus: 2N HC1, 0.1N HC1, 0.95N HC1, and N/64 HC1. Other terms are used in expressing the concentrations of solu- tions when reference is not made to a definite amount of solvent or solute; thus a dilute solution contains a relatively small amount of the solute as compared with that of the solvent; a concentrated 230 INORGANIC CHEMISTRY FOR COLLEGES solution contains a relatively large amount of the dissolved sul> stance. A saturated solution contains as much of the dissolved sub- stance as can be held in solution when some of the solid substance is present. The restriction stated in the last clause is necessary, because it is possible for a liquid to hold in solution more of a solid than can be presented in a saturated solution. This will be clear from the following considerations. The solubility of most salts increases with rise in temperature. If a saturated solution of such a salt is made at a certain temperature, and is then separated from any solid present, it is often possible to lower the tempera- ture and have all the salt remain in solution. Such a solution is said to be super-saturated, for if a bit of the solid is added, crystalli- zation takes place at once and the concentration of the solution decreases until it reaches that of a saturated solution when the solution and the solid are in equilibrium. 258. The Extent of the lonization of Acids, Bases, and Salts. It has been stated that the characteristic properties of solutions of acids are due to the hydrogen ions which are present. If this is true the behavior of all acids should be in general alike, but there should be a difference in activity and the rate at which they react, if the acids differ markedly in the extent to which they ionize. We find this to be the case. If we place pieces of zinc of equal size and shape in normal solutions of hydrochloric, sulphuric, and acetic acids, we find that the rate at which hydrogen is evolved in the three cases is different. A determination of the extent to which these acids ionize, by the methods outlined above, shows that they differ markedly in this respect. The extent to which an electrolyte ionizes has a marked effect on its chemical behavior. Acids which are largely ionized are often spoken of as strong acids because their characteristic acidic proper- ties are marked; acids that ionize but slightly are weak acids. For example, hydrochloric acid in a one-tenth normal solution at 18 is ionized to the extent of 92 per cent. Under the same con- ditions of concentration and temperature sulphuric acid is 61 per cent ionized, and nitric acid 92 per cent. These are all strong acids. Carbonic acid in one-twenty-fifth normal solution is ionized to the extent of 0.2 per cent only; it is a very weak acid. The hydroxides of sodium and potassium are strong bases; they are 91 per cent ACIDS, BASES, SALTS. SOLUTIONS 231 ionized in 0.1N solutions. Calcium and barium hydroxides are also strong bases; in N/64 solutions they are about 91 per cent ionized. Ammonium hydroxide is a relatively weak base; it is ionized to the extent of 1.3 per cent in 0.1N solutions. Most salts are largely ionized. The percentages of ionization of a few salts in 0.1N solutions at 18 are as follows: Sodium chloride 84, sodium nitrate 83. sodium sulphate 70, sodium acetate 79. The values for the potassium salts are about the same as those of the sodium salts. Copper sulphate and zinc sulphate are ionized to a less extent about 40 per cent and are typical of salts which yield bivalent ions. See also 599, page 511. EXERCISES 1. If you were given a sample of an element what experiments would you perform to determine whether it was an acid-forming or a base-forming element? 2. A sample of hydrochloric acid was neutralized with a solution of sodium hydroxide which contained 30 grams of the base in 1 liter of solution. In three experiments 16.42, 16.50, and 16.35 c.c. of the base were used to neu- tralize 10 c.c. of the acid. Calculate the weight of hydrogen chloride in 1000 c.c. of the solution of the acid. 3. Write the formulas of the ions formed when compounds of the follow- ing composition are dissolved in water: (a) BaCl 2 , (6) KNO 3 , (c) NiSO 4 , (d} Na,HP0 4 , (e) CoSO 4 , (/) A1C1 3 , (0) K 2 SO 4 , (/*) NaHCO 3 , (t) K 2 CO 3 , 0") Ca(N0 3 ) 2 , (/c) K 2 Cr0 4 , (0 NaHSO 3 . 4. The compounds having the composition represented by the following formulas are insoluble: (a) BaSO 4 , (6) CaCO 3 , (c) Ca 3 (PO 4 ) 2 , (d) AgBr, (e) CuS, (/) A1(OH) 3 , (g) Cu(OH) 2 . Write equations, using ionic formulas for reactions, by which each compound can be prepared by double decompo- sition. 5. Write ionic equations for the following: (a) the formation of ferrous chloride, FeCl 2 , by the action of hydrochloric acid on iron; (6) the elec- trolysis of a solution of salt; (c) the action of iron on a solution of copper sulphate, (d) the neutralization of calcium hydroxide by nitric acid. 6. An experiment showed that when 3 grams of a certain non-electrolyte were dissolved in 100 c.c. of water the boiling-point of the solution was 100.26. Calculate (a) the molecular weight of the compound and (6) the freezing- point of the solution. 7. A solution of 1.5 grams of a non-electrolyte in 50 c.c. of water was found to freeze at 0.6. Calculate (a) the molecular weight of the compound and (6) the boiling-point of the solution. 8. When 0.1 gram-molecular-weight of an acid that formed two ions was dissolved in 1000 c.c. of water, the freezing-point of the solution was found to be 0.357. Calculate the percentage dissociation of the acid. 232 INORGANIC CHEMISTRY FOR COLLEGES 9. What weights of the following must be dissolved in enough water to make 1000 c.c. of solution in preparing molar solutions (a) HC1, (6) H 2 SO 4 , (c) KOH, (d) Ba(OH) 2 , (e) BaCl 2 , (/) CuSO 4 ,5H 2 O? (g) What is the nor- mality of the solutions of the acids and bases? 10. Express the normality of solutions that contain (a) 42 grams HC1 per liter, (6) 23 grams H 2 SO 4 per liter, (c) 10 grams Ba(OH) 2 per liter. 11. It was found that 1.6 c.c. of a solution of hydrochloric acid were required to neutralize 1 c.c. of a normal solution of sodium hydroxide. What is the normality of the acid solution? CHAPTER XVIII CHEMICAL EQUILIBRIUM 259. Chemical reactions take place between molecules, and these are physical entities; as a consequence, for a complete inter- pretation of such reactions we must take into account not only the substances involved and their chemical properties, but the physical conditions under which the molecules are brought together, such as concentration and temperature. If we limit the dis- cussion for the present to gases, it is clear that since chemical reaction can take place only when the molecules approach one another and come into what might be called chemical contact, the concentration of the molecules is an important factor in the rate at which they interact. If in a given volume we have a given number of two kinds of molecules which react with each other at a certain rate, it is evident that if we introduce into this volume a larger number of these molecules the number of con- tacts between them in a given time will increase, and the rate of the interaction will be greater more molecules will be converted into the product of the reaction. Again, if we raise the temperature of the molecules they will move more rapidly, their contacts will be more frequent, as a consequence, and the rate of the reaction will increase. We see, therefore, that concentration and temperature physical conditions are important factors in determining the rate at which molecules interact. Up to the present the chemical aspects of the reactions studied have been emphasized, and attention has been centered on the products formed as the result of these reactions; but we are interested also in how these transformations come about the mechanics of molecular action in chemical change. The detailed study of chemistry from this point of view has led to generalizations of great value, which have not only broadened markedly our knowledge of chemical phenomena, but, as we shall 233 234 INORGANIC CHEMISTRY FOR COLLEGES see, have been utilized in solving many problems of the greatest technical importance. 260. Reversible Reactions. Many chemical reactions are reversible, that is, the substances formed as the result of the reac- tion are able, under the proper circumstances, to interact and pro- duce the substances from which they were formed. This can be expressed by "the following equation : where A and B represent elements or compounds which through the expenditure of chemical energy pass into the substances rep- resented by C and D. If the latter can react to form A and B the reaction is said to be reversible. 261. Chemical Equilibrium. It is evident that when a reaction of the above type takes place all four substances will be present. A and B react at a definite rate which is determined by their mutual chemical affinities; likewise C and D, but the rate in this case is different. When the reaction starts the amounts of A and B will decrease and those of C and D increase, but a time will come, as the reaction proceeds, when the quantity of any one substance formed just equals the quantity of it which undergoes trans- formation. When this condition has been reached the reaction is said to be in equilibrium. It is important to emphazise the fact that chemists are of the opinion that the molecules are in a state of constant change the equilibrium is dynamic, not static. This type of equilibrium has been emphasized already in con- nection with vaporization (179) . When such a state of equilibrium exists as the result of the continued change of molecules into other molecules of a different kind it is defined as chemical equilibrium. An example of a reaction which reaches an equilibrium has already been given (125): C1 2 + H 2 O <= HOC1 + HC1 In this case equilibrium results in a solution of 1 volume of chlorine in 1 of water when 66 per cent of the chlorine is present as molecules of the element and 34 per cent as hypochlorous and hydrochloric acids. 262. Effect of Concentration on Equilibrium. The concen- tration of the substances in a chemical reaction at equilibrium can CHEMICAL EQUILIBRIUM 235 be varied by adding more of one of the substances or by taking away some of one. In a typical reaction of this type, A + B<=C + D if more of A is added and the volume kept constant so that the con- centration of A is increased, there will be a larger number of con- tacts between A and B in a given time; as a consequence, the rate at which C and D are produced will be greater, and their amounts present when equilibrium is reached will be increased. It follows, therefore, as a consequence of increasing the concentration of one or more of the substances taking part in a reversible reaction, that a change takes place as a result of which more of the products pro- duced from these substances are present when equilibrium is attained. This fact expressed more definitely and in a mathe- matical form is known as the law of molecular concentration, which is sometimes called the law of mass action (see 264). An exam- ple will serve to emphasize this important conclusion. At high temperatures the reaction expressed by the following equation is reversible : 2SO 2 + O 2 <= 2SO 3 Sulphur dioxide reacts with oxygen to form sulphur trioxide, but at the same temperature the trioxide decomposes into sulphur dioxide and oxygen. It has been found by experiment that when the dioxide and oxygen are mixed in the proportion represented by the equation and heated to 720 about 40 per cent of sulphur di- oxide is converted into trioxide. If, however, twice as much oxygen is used and the volume of the gases kept the same as before so that the concentration of the oxygen is doubled, about 60 per cent sulphur dioxide is changed. The important consequences of the change in the equilibrium with change in concentration are evident. If we are preparing a substance by means of a reaction which is reversible and this is often necessary we can increase the amount of product formed, by increasing the concentration of one or more of the reacting sub- stances. The excess of the material which is present when equi- librium is reached can be recovered; its presence during the reac- tion was important since it effected the result attained, but it was not used up. 236 INORGANIC CHEMISTRY FOR COLLEGES The concentrations of the substances taking part in a rever- sible reaction can be changed by taking away one of the products of the reaction. If, for example, in the reaction represented by the equation C is removed as fast as it is formed, it is evident that there is no opportunity for C and D to produce A and B. As a consequence, the reaction runs to completion and A and B are completely trans- formed into C and D. An important example of this kind of a reaction is what we have called double decompositions (148). When, for example, sodium chloride is treated with sulphuric acid the reaction represented by the following equation takes place : NaCl + H 2 SO 4 = NaHSO 4 + HC1 In this case the hydrogen chloride is a gas and escapes and cannot, therefore, interact with the sodium hydrogen sulphate to form sodium chloride. The reaction is a reversible one, however; if it is carried out in such a way that the hydrogen chloride cannot escape, all of the salt is not changed to sodium hydrogen sulphate. In fact, by selecting the proper physical conditions, we can con- vert sodium hydrogen sulphate and hydrogen chloride into sodium chloride and sulphuric acid and thus completely reverse the reac- tion. This can be done by adding the sulphate to a concentrated solution of hydrochloric acid. In this case the hydrogen chloride is kept in the field of reaction by being dissolved in water. Sodium chloride is insoluble in concentrated hydrochloric acid and pre- cipitates. The reaction in this case is expressed by the above equation written backward: NaHSO 4 + HC1 = NaCl + H 2 SO 4 A double decomposition takes place, since the sodium chloride is automatically removed as the result of the fact that it is insoluble in the liquid present. The high concentration of the hydrogen chloride in the saturated solution is a factor in forcing the equi- librium to completion as indicated by the equation. From the standpoint of chemical equilibrium we can now under- stand the reason for the statement which has been given as to the conditions under which a reversible reaction becomes one of double CHEMICAL EQUILIBRIUM 237 decomposition. This occurs, it will be recalled, when, as the prod- uct of the reaction, there is formed a gas, a substance which decom- poses into a gas, or an insoluble substance. In these cases the product which possesses any one of the above properties is auto- matically removed from the sphere of action. Double decom- positions which result from the formation of undissociated substances are explained as follows : In solution, most inorganic reactions take place between ions, and these play the same part as the molecules do when reactions occur between the latter. For example, in a reaction of neu- tralization, such as that between sodium hydroxide and hydro- chloric acid, the theories accepted at present lead to the view that the reaction takes place between ions which are represented in the following equation: Na + + OH~ + H + + Cr = Na + + Cl" + H 2 O Since the sodium ion and the chlorine ion do not take part in the reaction, the latter can be simplified to the following : H + + OH~ => H 2 O The equilibrium in this reaction is such that practically all of the ions change into undissociated water; as the result, the action of an acid and a base is one of double decomposition. 263. Equilibrium in Homogeneous and in Heterogeneous Sys- tems. In the discussion up to this point it has been assumed that all of the substances involved in the reactions are present in one phase (188) that is, they are uniformly distributed. In this case the system is said to be homogeneous. If two or more of the sub- stances are present in different phases, for example, if one is a gas and one is a solid, the system is said to be heterogeneous. If we are considering an equilibrium in the gaseous or liquid phase in such a system, the concentration of the solid in these phases is constant as long as any of the solid is present. For example, in the reversible reaction represented by the equation 4H 2 O + 3Fe * Fe 3 O 4 + 4H 2 the relation between the amounts of water and hydrogen present at equilibrium is not changed by altering the amounts of iron or iron oxide, provided some of each is present. 238 INORGANIC CHEMISTRY FOR COLLEGES 264. The Law of Molecular Concentration. The effect of concentration on chemical equilibrium can be put into a simple mathematical form which makes it possible to calculate how much of each of the reacting substances is present at equilibrium at any concentration, provided the amounts present at one concentration are known. It is necessary to find these amounts for one set of concentrations by experiment. In the reversible reaction repre- sented as follows : the rate at which A and B are transformed into C and D is pro- portional to the molecular concentration of A and B and to the mutual chemical affinities between the latter. Likewise, the rate at which C and D are transformed into A and B is proportional to the chemical affinity which exists between C and D and to their molecular concentrations. Since the chemical affinities are constant at a fixed temperature, the rates of the two reactions under these conditions are proportional to the concentrations only. It will be recalled (note, section 88) that if a quantity varies as two other quantities vary, the first will be proportional to the product of the latter; and that if one quantity is proportional to another, the first is equal to a constant multiplied by the second. Making use of these mathematical modes of expression, we can put the statement in regard to the effect of concentration on equilibrium into the form of a formula. The rate, R' of the reaction A + B =C + Dis proportional to the molecular concentration of A,(C A ) and of B 3 (C B ); if K' represents a constant, then R' = K' X C A X C B The similar equation for the opposed reaction C-r-E) =A + Bis R" = K" X C c X C D At equilibrium these rates are equal, and, therefore, under these circumstances R' = R" and K' X C A X CB = K" X C c X C D , or C>C X OD -K- -rr- C A XC B = K 77 = K CHEMICAL EQUILIBRIUM 239 It is customary in writing equations to express equilibria, to put in the numerator of the fraction the molecular concentrations of the products represented on the right side of the chemical equation for the reaction involved. The last equation given above states that in a reversible reaction the product of the molecular concentra- tions of the substances formed, divided by the product of the molec- ular concentrations of the interacting substances, is a constant. This constant is known as the equilibrium constant of the reaction. The general statement given above is one mode of expression of the so-called law of molecular concentration. It was proposed in 1864 by Guldberg and Waage, two Norwegian chem- ists, and was based on their own experiments as well as those of Wilhelmy, who had studied the influence of concentration on the rate of reactions, and those of Gladstone, who had investigated the influence of concentration on chemical equilibrium. The law was originally called the mass law but the name was not well chosen, since it does not apply to mass but to molecular concentration. For example, in using the law in connection with a particular re- action we do not express the concentration in grams, by which mass is measured, but in moles; the law was deduced as the result of the consideration of the effect on the equilibrium of an in- crease in the number of molecules taking part in any reaction. A knowledge of the value of the equilibrium constant is of the greatest importance when a reversible reaction is being used for the preparation of any of the substances involved, for it makes it possible to calculate to what extent changes in concentration will effect the quantity of the desired substance obtained. If a reversible reaction is of the type represented by the equa- tion A + 2B = C + D, the mathematical expression of the law takes a slightly different form. In this case two molecules of B react with one of A, or A + B + B=C + D; accordingly, R' = K' X C A X C B X C B or R' = K' CA X C B 2 . It is seen from this that the most gen- eral expression of the law is as follows : Cp U X Cp v . . . _ -p,- C A " X CB* = 240 INORGANIC CHEMISTRY FOR COLLEGES where u, v, w, and x represent, respectively, the number of mole- cules of the several substances which enter into the reaction. 265. Effect of Temperature on the Rate of Chemical Reac- tions. Since the equilibrium attained in a reversible reaction is dependent on the relative rates at which the two opposing reac- tions take place, it is necessary to consider the effect of change in temperature on these rates. It has been found by experiment that within ordinary ranges of temperature the rate at which reac- tions take place is approximately doubled for a rise of 10 degrees. From the standpoint of the kinetic theory this fact appears reason- able, for with rise in temperature it is postulated that the molecules move faster and come into contact more frequently; as a conse- quence, more of the products of the reaction are formed in a given time the rate of the reaction increases. The fact that reactions proceed more rapidly as the tempera- ture is increased is utilized by the chemist. A reaction proceeding at a certain rate at room temperature (20) will go approximately 256 times as fast at 100. The change in temperature in this case is 80 degrees, and since the rate is doubled for each 10 degrees, the rate is 2 8 = 256 times as fast at the higher temperature. Many inorganic reactions, such as double decompositions, take place practically instantaneously and heating is unnecessary. When reactions take place slowly, however, such as the decomposition of potassium chlorate into potassium chloride and oxygen, they are usually carried out at more or less elevated temperatures. 266. Effect of Temperature on Chemical Equilibrium. When a system in chemical equilibrium is heated, the rates of the two opposing reactions are both increased. This change results from the increased motion of the molecules and the change in the chem- ical affinities between the interacting substances. If the tempera- ture of any definite system in equilibrium is raised, the effect of the increase is the same on the motion of all the molecules, and the change in the rates of the opposing reactions due to this cause alone will be the same. But increase in temperature changes the inten- sity of the chemical affinities of all molecules differently; conse- quently the change in the rates at which the two reactions proceed will be different when the temperature of the system is raised, and the concentrations at equilibrium will, as a result, be different. The effect of temperature on chemical equilibrium has been CHEMICAL EQUILIBRIUM 241 extensively investigated on account of its importance. It was found to be associated, as one might expect, with the amounts of heat evolved or absorbed when chemical reactions occur. If we take, for example, the reaction represented by the following equation, A + B<=C-f-D + z cal. when A and B react to form C and D, heat to the amount of x calories is set free; on the other hand when C and D react to form A and B the same amount of heat is absorbed transformed into chemical energy. The change of heat energy into chemical energy takes place more readily at high than at low temperatures; as a consequence, when the temperature of a system represented by the equation given above is raised, the rate of the reaction by which C and D change to A and B is increased and the equilibrium shifts so that more of the latter are present. The equilibrium is shifted from right to left in the above equation by increasing the concen- tration of C or of D or the temperature at which the reaction is car- ried out. It is evident that if we are using such a reaction for the preparation of C or D, it should be carried out at the lowest tem- perature possible, at which the rate is satisfactory. Since the rate of the reaction may be very small at the temperature which yields a satisfactory concentration of the desired substance, such reac- tions are often brought about in the presence of a catalyst. It will be recalled that catalysts markedly increase the rate of reac- tions. In the case of an endothermic reaction such as, E + F G + H-z cal. rise in temperature leads to the shifting of the equilibrium from left to right. In this case when E and F change into G and H heat is absorbed, and as heat can be more readily changed into chemical energy at high temperatures, rise in temperature favors the forma- tion of G and H. The direction in which the equilibrium shifts can also be seen by changing the equation to read, E + F + x cal. <= G + H and using the method employed in the previous example; it is seen that increase in temperature shifts the equilibrium so that more G and H are formed. 242 INORGANIC CHEMISTRY FOR COLLEGES It follows from the above that if a substance is to be prepared through the use of an endothermic reaction, the preparation should be carried out at as high a temperature as possible. Under these conditions the rate of the reaction is usually high and no catalyst is required. In general, when a large amount of heat is evolved or absorbed in a reaction, the effect of change in temperature on the equilibrium is great; on the other hand, when the heat change is small the effect of change in temperature is small. A knowledge of this principle is of great value in coming to a conclusion as to whether a par- ticular reaction could be used for the preparation on the industrial scale of one of the products of the reaction. For example, if the reaction is an endothermic one and the equilibrium at a definite temperature is such that but a small proportion of the desired product is formed, whether or not the reaction can be used will depend, among other considerations, on the heat exchange. If this is great, the proportion of the desired product at high tempera- tures will be much increased enough perhaps to make it advisable to use the reaction. If the heat change is small, increase in tem- perature will have little effect on the equilibrium. 267. The Law of van't Hoff. The effect of temperature on equilibrium was studied by van't Hoff, a Dutch chemist, who generalized the facts in regard to it into a form which is known as van't Hoff's law of mobile equilibrium. This law states that when the temperature of a system in equilibrium is raised, the equilibrium-point is displaced in the direction which absorbs heat. This law is of the greatest importance and along with the law of molecular concentration is the guiding principle followed in the study and utilization of all reversible reactions. The law of van't Hoff applies to physical equilibria, such, for example, as that existing between a salt and its saturated solution. If we represent this equilibrium as follows: Solid salt + water ^ solution of salt x cal. the fact is indicated that when the salt dissolves, heat is absorbed. Applying van't Hoff's law we would expect that with rise in tem- perature the solubility would increase and this is the fact. Anhy- drous sodium sulphate is an example of another type of salt; when it dissolves heat is given off. In this case the solubility decreases with rise in temperature. CHEMICAL EQUILIBRIUM 243 When sodium chloride, common salt, dissolves in water the heat change is small ( 1,180 cal.), and the effect of rise in tem- perature on its solubility is small. On the other hand, when potas- sium nitrate dissolves, a much larger amount of heat is absorbed ( 8,520 cal. per mol) ; this results in the fact that the temperature of the solution drops when a strong solution of the salt is made. In this case the solubility of the salt is markedly affected by rise in temperature. 268. The Law of Le Chatelier. The law of van't Hoff applies to certain cases coming under the broader generalization first put forward by Le Chatelier. This law, which applies to both physical and chemical equilibria, sums up the behavior of substances when they are subjected to a change in conditions. It is as follows: If some stress is brought to bear on a system in equilibrium, the equilibrium is displaced in such a direction that the normal effect of the stress results. 1 By stress is meant the effect of any thing or any kind of energy which has an effect on the equilibrium; it may be, for example, change in concentration, pressure, or tem- perature. 269. The Effect of Pressure on Equilibrium. When a chemical reaction takes place which is accompanied by a change in volume, the pressure under which the change occurs has a marked effect on the equilibrium. If we consider the reaction by which ammonia is formed from hydrogen and nitrogen, N 2 + 3H 2 ^ 2NH 3 we see that 1 volume of nitrogen unites with 3 volumes of hydrogen to form 2 volumes of ammonia gas there is a contraction from 4 to 2 volumes. By making use of Le Chatelier's law we can discover in which direction the equilibrium will shift if the pressure on the system is increased. If we replace the words " some stress " in the above statement of the law by the word pressure, the law will read, if pressure is brought to bear on a system in equilibrium the equilibrium is displaced in such a direction that the normal effect of the pressure results. Increase in pressure brings about a decrease in volume; as a consequence, in the above .reaction under increased pressure more ammonia will be present in the system at 1 This law is usually expressed as follows : If some stress is brought to bear on a system in equilibrium, the equilibrium is displaced in the direction which tends to undo the effect of the stress. 244 INORGANIC CHEMISTRY FOR COLLEGES equilibrium because when the change of nitrogen and hydrogen to ammonia takes place the volume decreases. This particular reaction will be considered in detail later (339), and we shall see that a knowledge of the facts summarized in Le Chatelier's law made it possible to develop a process for the manufacture of am- monia. EXERCISES 1. Salts having the following formulas are practically insoluble in water: (a) Agl, (6) BaCO 3 , (c) CaCO 3 , (d) CuS. Write equations for reactions by which each can be formed, and state how the principle of chemical equi- librium applies to the reactions. 2. Zinc sulphide, ZnS, is slightly soluble in water, but its solubility increases in the presence of hydrochloric acid, the amount passing into solu- tion being determined by the concentration of the acid. State reasons for what you would expect to happen if a solution of (a) Na 2 S and one of (b) H 2 S were added to a solution of ZnCl 2 . Copper sulphide, CuS, is insoluble in hydrochloric acid, (c) What would happen if solutions a and b above were added to separate portions of a solution of copper chloride, CuCl 2 ? 3. Nitric acid boils at 86. Under what conditions could it be prepared by double decomposition from sodium nitrate, NaNO 3 , and sulphuric acid? 4. Discuss the changes which occur in the following systems when heat is applied: (a) H 2 O (liquid) + H 2 O (gas); (b) H 2 O (solid) <= H 2 O (liquid); (c) Discuss the changes which occur in the two cases when the temperature is constant but the pressure is increased. 5. Discuss the changes produced in the reversible reaction PC1 3 -f C1 3 =^ PCU + 33,000 cal. when (a) the temperature is changed and (b) the pressure is changed, (c) If PC1 6 is converted into vapor in a space filled with C1 2 how would its dissociation compare in extent with that produced when it is vaporized in the same space filled with air? 6. Show how Henry's law is consistent with LeChatelier's law. 7. Would you expect pressure to have much effect on the solubility of a salt in water? Give a reason for your answer. Why does the temperature have a marked effect on the solubility of salts? 8. Lead iodide, which is slightly soluble in water, can be prepared by the following reaction: Pb(NO 3 ) 2 + 2KI = PbI 2 + 2KNO 3 . (a) Write the reaction, using ionic symbols, and state what would happen if a small amount of a solution of lead nitrate were added to a saturated aqueous solution of lead iodide? (b) How could the same result be produced without the addi- tion of a lead salt? 9. Silver chloride, AgCl, which is slightly soluble in water, is frequently precipitated and weighed in making a quantitative determination of chlorides. When it is washed to remove impurities a small amount of silver nitrate is added to the wash-water. State a reason for the use of the latter. CHAPTER XIX SULPHUR AND HYDROGEN SULPHIDE 270. Sulphur occurs in the free condition in volcanic regions and has been known since the earliest times. It is mentioned in the Bible and in Homer, and was considered to be one of the ele- mentary principles of which the earth is composed. It represented the spirit of fire and was called brimstone (fire-stone). In the fifteenth century when the action of chemical substances on the body was first studied in an endeavor to find some way to combat disease, sulphur was used as a medicine. It was believed that the body was made up of the elementary principles mercury, sul- phur, and salt, and a lack in one of these produced illness. In the case of certain ailments, sulphur was prescribed and in the case of others, compounds of mercury or various kinds of salts. As the result of the foundation of a system of medicine on this hypothesis, the effect on the body of a great many chemical substances was discovered, and a basis was laid for that branch of modern medicine which is called pharmacology. The utilization of sulphur as a medicine has extended to modern times, although at present its application is limited to use in ointments and salves to combat certain diseases of the skin. Sulphur is a necessary constituent of the body, and it must be present in our food, but it has been shown that the free element is not assimilated; it must be present in com- plicated organic compounds to be of value, and it is from these, such as the protein of eggs, meat, etc., that we obtain our supply of the element. Any excess of sulphur over that required is excreted as sulphates in the urine. 271. Occurrence of Sulphur. Free sulphur is found abundantly in Sicily, mixed with limestone, CaCOa, gypsum, CaSCU, 2H 2 O and other minerals. It also occurs in Japan, Italy, Spain, Cali- fornia, Egypt, China, and India. Sulphur is found in rock deposits where it has been probably formed from gypsum through the action 245 246 INORGANIC CHEMISTRY FOR COLLEGES of bacteria. Large quantities of sulphur are found in the sedi- mentary deposits of Texas and Louisiana. The element also occurs in compounds; hydrogen sulphide, H2S, which is a gas with a disagreeable and characteristic odor, issues from volcanoes, and occurs dissolved in the water of sulphur springs; sulphur dioxide, SOo, is also a product of volcanic activity. Many sulphides are important minerals, some of which are used as a source of sulphur and some of the other elements which they con- tain; pyrite, FeS2, furnishes sulphur for the manufacture of sul- phuricacid ; zincblende, ZnS, cinnabar, HgS, galena, PbS, and chalco- cite, Cu2S, are valuable ores. Many sulphates occur in nature, of which gypsum, CaSC>4, 2H2O, is perhaps the most important. Sulphur occurs in all living things; it is a constituent of certain proteins and is found in hair, nails, meat, eggs, etc. When these compounds undergo putrefaction the sulphur which they contain is converted into hydrogen sulphide and other substances which pos- sess a disagreeable odor. The element occurs in many oils found in plants, which give the latter their most characteristic properties ; such oils have been extracted from onions, mustard, garlic, cab- bage, etc. 272. Extraction of Sulphur. The chief sources of sulphur are Sicily and Louisiana. In Sicily the rock with which the sulphur is mixed is piled in a crude kiln and covered with some of the ore left from a previous burning. The kiln is fired by igniting some of the sulphur, and allowing it to burn for some time; the draught holes are then closed. The heat generated melts the sulphur, which collects in the bottom of the kiln and is drawn off from time to time and run into wooden molds. This method is crude and wasteful; about one-quarter of the sulphur burns to sulphur dioxide, which passes into the air, and as a result damages vegetation in the neighborhood of the kiln. The sulphur in Louisiana, which is found in deposits over 100 feet thick at a depth of about 450 feet, is obtained by an ingenious method devised by Frasch. Four concentric pipes, 1, 4, 6, and 10 inches in diameter were sunk into the deposit. Between the 6- and 10-inch pipes superheated water at about 170 is pumped into the sulphur, which melts at this temperature. Hot air under pressure is forced down the 1-inch pipe, and, as a result, a mix- ture of air, molten sulphur, and water rises in the space between the SULPHUR AND HYDROGEN SULPHIDE 247 4-inch and 6-inch pipes. The material is run into wooden tanks where the sulphur solidifies. The product is quite pure and is used without further refining. Crude sulphur is refined by distilling it from iron retorts and conducting the vapor produced into chambers made of brick. If the temperature of the condensing chamber is below 110 the sul- phur collects in the form of a fine powder, called " flowers of sul- phur"; at a higher temperature the vapor condenses to a liquid, which is run off into cylindrical molds, where it solidifies. In this form the sulphur is often called roll sulphur, or roll brimstone. 273. Physical Properties of Sulphur. It will be recalled that carbon can exist in three distinct allotropic forms diamond, graphite, and amorphous carbon. Sulphur shows the property of existing in two crystalline allotropic modifications, but in this case but one form is stable at room temperature and the other slowly changes to it on standing. The stable form of sulphur, which occurs native, crystallizes in the rhombic system and is called, therefore, rhombic sulphur. It is pale yellow in color, is brittle, odorless, tasteless, has the specific gravity 2.06 and melts at 114.5; it is almost insoluble in water, but dissolves in carbon disulphide (41 parts in 100 at 18), forming a solution from which it can be obtained in well-defined crystals by evaporation. When rhombic sulphur is melted and allowed to cool, the crystals formed are long, transparent needles which belong to the monoclinic system. These can be obtained by melting some sul- phur in a crucible, and before it has completely solidified, pouring off the liquid through a hole made in the solid crust formed on the surface of the liquid. Monoclinic sulphur is almost colorless, melts at 119.25, has the specific gravity 1.96, and dissolves in carbon disulphide. It slowly changes at room temperature to rhombic sulphur, and the transparent needles become opaque as the result of the formation of minute rhombic crystals. The two solid forms of sulphur are in equilibrium at 96, which is the transition point for the two forms. Below 96 rhombic sulphur is the stable form. When either form of sulphur is melted, a clear, pale-yellow, limpid liquid is first obtained; as the temperature is raised the color changes to dark brown and the liquid becomes at about 160 so viscous that it will scarcely flow out of the vessel containing it, when it is inverted. As the temperature is increased above 260 248 INORGANIC CHEMISTRY FOR COLLEGES the mixture becomes less viscous, and at 444.7 the liquid boils. The yellow liquid is represented by the symbol SX and the brown liquid by Sju. If S/z is allowed to cool slowly it passes into SX and the crystals obtained on solidification dissolve in carbon disul- phide. If, however, the brown molten sulphur is poured into water and chilled suddenly in this way, the sulphur assumes the form of a plastic mass which is somewhat elastic. After standing some days plastic sulphur becomes hard and opaque; it consists of a mixture of rhombic sulphur and S/x in the form of an amorphous solid; the two forms can be separated by carbon disulphide in which S/i is insoluble. SM changes very slowly to rhombic sulphur at room temperature, a number of years being necessary to effect complete transformation. 274. Chemical Properties of Sulphur. Sulphur is an active element and unites with both metallic and non-metallic elements. In this respect it resembles oxygen and many of the formulas of the sulphides are analogous to those of the oxides. This is due to the fact that when sulphur unites with metals it has the valence 2, which is the valence of oxygen. The formulas of some com- pounds produced as the result of the direct union of sulphur with other elements are as follows: H 2 S, ZnS, CuS, FeS, Ag 2 S, P 2 S 5 , and 082; the formulas suggest those of the oxides of these ele- ments. In compounds containing the more negative elements sulphur exhibits the valence 4 or 6; thus, when it burns sulphur dioxide, SO 2 , is formed along with a small amount of sulphur tri- oxide, SOa; with chlorine, sulphur forms compounds of the com- position S 2 C1 2 , SC1 2 , and SCU. Sulphur monochloride, S 2 C1 2 , is a reddish-yellow liquid, boiling at 138, which dissolves sulphur; the solution is used in vulcanizing rubber. Sulphur combines slowly at room temperature with all metals except the least active ones like platinum and gold. Even silver which is not affected by oxygen reacts readily with sulphur. If a bit of rubber is left in contact with silver the latter turns black as the result of a reac- tion between the metal and the sulphur present in the vulcanized rubber. Silver spoons are tarnished when left in contact with eggs, because the latter contain sulphur compounds which are decomposed as the result of the affinity of silver for sulphur. Many metals burn when introduced into sulphur vapor. The union of iron and sulphur has already been described (16). SULPHUR AND HYDROGEN SULPHIDE 249 Carbon combines with sulphur, but as the reaction is an endo- thermic one, energy must be supplied to bring about the formation of carbon disulphide (215). 275. Uses of Sulphur. In the crude condition sulphur is used for making sulphur dioxide, which, in turn, is employed in the manufacture of sulphuric acid, in bleaching, and for many other purposes. Large quantities of flowers of sulphur are consumed in destroying a fungus which causes a disease in grapevines; the value of the element for this purpose is probably due to the fact that in the presence of the oxygen and water in the air it slowly changes to sulphuric acid, which prevents the growth of the fungus. Refined sulphur is used in the manufacture of matches, gunpowder, fire- works, and for vulcanizing rubber. Many of the important newer dyes are made by heating with free sulphur certain compounds prepared from the products found in coal-tar; some of these dyes contain traces of free sulphur or compounds which give up the element readily, and, as a consequence, blacken silver which is wrapped in material dyed with them. The industrial importance of sulphur can be seen from the fact that over 800,000 tons of it are used annually. HYDROGEN SULPHIDE 276. Hydrogen sulphide, H2$, is found in the gases which issue from volcanoes and in the water of sulphur springs. It is formed in the putrefaction of certain proteins, which are the most impor- tant nitrogenous constituents of animal and vegetable substances. The disagreeable odor produced as the result of the decomposition is in part due to this gas. Rotten eggs are said to smell of hydrogen sulphide, but an egg can have undergone decomposition and possess a most disagreeable odor before a test will show the presence of the gas. 277. Preparation of Hydrogen Sulphide. Hydrogen sulphide can be made by passing hydrogen through boiling sulphur, but the method lacks practical significance: H2 H- S = H2S It is prepared in the laboratory by the action of hydrochloric acid on ferrous sulphide : FeS + 2HC1 = FeCl 2 + H 2 S 250 INORGANIC CHEMISTRY FOR COLLEGES The reaction is similar to that between ferrous oxide and hydro- chloric acid: FeO + 2HC1 = FeCl 2 + H 2 O It has been pointed out that sulphur resembles oxygen in many of its reactions; the behavior of sulphides in certain cases is similar to that of oxides. There are differences, however; every metal forms an oxide which will react with hydrochloric acid, but all sulphides do not dissolve in the acid. The sulphides of the metals from potassium to iron inclusive, in the electro-motive series of the elements (252), dissolve in dilute hydrochloric acid. Any of these could be used to make hydrogen sulphide, but iron sulphide is commonly employed because it is readily prepared from iron and sulphur. Other acids can replace hydrochloric acid, but nitric acid can not, because it oxidizes hydrogen sulphide. When many organic substances which contain hydrogen and sulphur are heated to a high temperature or are burned in an insufficient supply of air, hydrogen sulphide is formed. A decom- position of this kind leads to the formation of hydrogen sulphide when coal is heated to prepare coal-gas, or when it is burned with insufficient draft to produce complete combustion (223). 278. Physical Properties of Hydrogen Sulphide. Hydrogen sulphide is a colorless gas, somewhat heavier than air; 1 volume of water dissolves 4.37 volumes of hydrogen sulphide at 0; 3.60 at 10, and 3.23 at 15; the gas is insoluble in boiling water. On account of the solubility of hydrogen sulphide in water it is usually collected in the laboratory by upward displacement of air. The critical temperature of the gas is 100; below this it can be lique- fied by pressure. Hydrogen sulphide boils at 62 and melts at 83. The gas is a powerful poison; it produces nausea and coma and 1 part in 200 of air causes death. 279. Heat of Formation of Hydrogen Sulphide. When hydro- gen unites with sulphur to form hydrogen sulphide a comparatively small amount of energy is set free. The thermo-chemical equa- tion when solid sulphur is used is as follows: H 2 + S = H 2 S + 4,800 cals. In the formation of water much more energy is liberated; when 2 grams of hydrogen burn in oxygen 58,100 calories are liberated. SULPHUR AND HYDROGEN SULPHIDE 251 We should expect, therefore, to find hydrogen sulphide a much less stable compound than water, and being unstable, more react- ive. When a large amount of chemical energy is transformed into heat in the formation of a compound, the compound is usually stable when heated. 280. Chemical Behavior of Hydrogen Sulphide. With the above facts in mind we are not surprised to find that hydrogen sulphide begins to dissociate into hydrogen and sulphur at a com- paratively low temperature, 310, as compared with that at which water begins to dissociate, about 1800. Hydrogen sulphide burns in air with a pale blue flame : 2H 2 S + 3O 2 = 2H 2 O + 2SO 2 In this case a large amount of energy can be changed into heat as the result of the union of hydrogen and sulphur with oxygen; when 1 gram-molecular-weight of hydrogen sulphide is burned 122,500 calories are set free. In general, compounds which are combustible are composed of elements that burn. At room temperature a solution of hydrogen sulphide is decom- posed slowly by the oxygen in the air. Under these conditions the hydrogen is oxidized to water and sulphur separates as a white powder: 2H 2 S + 2 = 2H 2 + 2S A similar reaction takes place when hydrogen sulphide is burned in a limited supply of air. All the metals down to and including silver in the electromotive series of the metals decompose hydrogen sulphide in the cold, and hydrogen and the sulphide of the metal are formed. The tarnishing of metals in air containing hydrogen sul- phide is due to this cause. If an exceedingly thin layer of sulphide is formed on silver it has a yellow or golden tint, but as the amount of sulphide increases the color changes to black. Hydrogen sulphide reacts with oxygen not only when it is in the free condition but when it is combined with other elements. For example, it reacts with sulphur dioxide according to the fol- lowing equation: 2H 2 S + SO 2 = 2H 2 O + 3S We see again in this reaction the strong tendency of hydrogen to unite with oxygen. It is probable that the sulphur found near 252 INORGANIC CHEMISTRY FOR COLLEGES volcanoes is produced as the result of this reaction. When hydrogen sulphide reacts in this way it is said to be a reducing agent, for it takes oxygen away from the sulphur; as a result of the reaction the sulphur dioxide is reduced and the hydrogen sulphide is oxi- dized. A solution of hydrogen sulphide in water shows an acid reac- tion with litmus, but it is a very weak acid as it is dissociated in one-tenth normal solution to the extent of 0.07 per cent only. It ionizes in steps as other dibasic acids do: H 2 S -> H + + HS~ The ionization according to the second expression is very slight, being less than that of water. When hydrogen sulphide is passed into a solution of sodium hydroxide, sodium hydrogen sulphide, NaHS, is formed : Na + + OH~ + H + + HS~ = Na + + HS~ + H 2 O The solution of the salt is neutral on account of the fact that practically no hydrogen ions are present. 281. If a solution of sodium hydrogen sulphide is mixed with one of sodium hydroxide and the water is driven off by heat, the normal salt is obtained : NaOH + NaHS <=> Na 2 S + H 2 O The reaction is a reversible one; if sodium sulphide is dissolved in water the reaction indicated by reading the above equation from right to left takes place. Sodium hydroxide and sodium hydrogen sulphide are formed, and the solution shows a strongly alkaline reaction. In this case hydrolysis is said to have taken place. Many substances are hydrolyzed by water and reactions of this kind are of importance. The change which takes place when water reacts with a compound and converts it into two or more compounds is called hydrolysis. The action of an acid and a base to form a salt and water is called neutralization (239); the reverse of this, the action of water with a salt to form an acid and a base is an example of hydrolysis. In some cases, as in the above, an acid salt is formed. The cause of hydrolysis is clear from a considera- SULPHUR AND HYDROGEN SULPHIDE 253 tion of the ions involved in the reaction. The equation for the action of water on sodium sulphide is as follows : 2Na + + S~ + H + + OH~ = 2Na + + HS~ + OH~ Water breaks down to a very small extent into H + and OH~ ions, the H 4 " from the water and the S~~ from the sulphide unite to form HS~, and, as a result, OH ions are left in the solution which, accordingly, shows an alkaline reaction. See section 600, page 511. 282. Sulphides. All the metals form sulphides; the com- pounds differ in solubility so widely that one of the most impor- tant parts of qualitative chemical analysis involves the preparation and separation of sulphides. The behavior of solutions of the salts of the various metals when hydrogen sulphide is passed through them can be illustrated by a few typical examples. The gas is passed successively into solutions of the following salts, sodium chloride, calcium chloride, zinc sulphate, ferrous chloride, copper sulphate, and arsenic chloride. Nothing takes place in the first two cases; in the others, precipitates are produced as the result of the formation of the sulphides of the elements. The equations for the reactions are as follows: ZnSO 4 + H 2 S = ZnS + H 2 SO 4 FeCls + H 2 S = FeS + 2HC1 CuSO 4 + H 2 S = CuS + H 2 SO 4 2AsCl 3 + 3H 2 S = As 2 S 3 + 6HC1 The sulphide of zinc is white, that of arsenic yellow, while ferrous sulphide and copper sulphide are black; the color of the sulphides is, thus, helpful in their identification. If dilute hydrochloric acid is added to the four sulphides, only those of zinc and iron dissolve; these sulphides are not precipitated if the solution is acidified before hydrogen sulphide is introduced, because they are soluble in acids. The reactions indicated by the first and second equations above are reversible. Hydrogen sulphide does not completely precipitate these sulphides, because as a result of the reaction an acid is produced. If it is desired to have the reaction proceed to comple- tion, a salt of hydrogen sulphide must be used, and for this pur- pose ammonium sulphide is generally employed: ZnSO 4 + (NH 4 ) 2 S = ZnS + (NH 4 ) 2 SO 4 254 INORGANIC CHEMISTRY FOR COLLEGES Under these circumstances no acid is formed and the reaction runs to completion. The behavior of the salts of the metals illustrated above is typical of many others; some are not precipitated by hydrogen sulphide from an aqueous solution (potassium to aluminium inclusive in the electromotive series), some are not precipitated in the presence of acids but are precipitated by ammonium sulphide (manganese to iron inclusive), and some are precipitated by hydro- gen sulphide in the presence of dilute acids (cobalt to gold inclu- sive). Of the sulphides iasoluble in dilute acids those of tin, arsenic, antimony, platinum, and gold are soluble in yellow ammonium sulphide, (NH^S,. By means of the reactions outlined above the metals can be separated into various groups (see Appendix VI) and by making use of the action of other reagents the metals in a single group can be separated and identified. These tests will be given later when the metals are described separately. 283. Test for Sulphides. Hydrogen sulphide is recognized by its characteristic odor or by the fact that it produces a coloration when it comes into contact with a piece of paper moistened with a solution of a lead salt; lead acetate is usually used in making the test. The equation for the reaction is H 2 S + Pb(C 2 H 3 O2)2 = PbS + 2H(C 2 H 3 O 2 ) The color of the lead sulphide precipitated varies from yellow to black according to the amount of gas present. In testing for a sulphide the substance is warmed with dilute hydrochloric acid, and if a gas is given off it is tested by exposing it to a piece of moist lead acetate paper. Many sulphides do not react with dilute hydrochloric acid. In testing for these a small piece of zinc is added to the mixture of acid and the substance under examination. If a sulphide insoluble in acids is present it is reduced by the nascent hydrogen, and hydrogen sulphide is formed. The equations for the reactions in the case of copper sulphide are as follows: Zn + 2HC1 = ZnCl 2 + 2H CuS + 2H = Cu + H 2 S The gas given off is tested as before with lead acetate paper. SULPHUR AND HYDROGEN SULPHIDE 255 284. Polysulphides. Sulphur dissolves in solutions of sodium sulphide and forms a mixture of compounds to which the formulas Na22, Na2$3, and Na2Ss have been assigned. If the solution containing these salts is poured into a dilute acicj, an oil is obtained from which liquid sulphides of hydrogen have been isolated. The one having the formula 11282 is unstable and resembles somewhat hydrogen peroxide, H2O2, in properties. Of compounds of this type, which are known as polysulphides, the one prepared by dis- solving sulphur in ammonium sulphide is used in qualitative analy- sis. Its formula is usually written (NH^S*, the x indicating that the substance is a mixture of polysulphides. Sodium poly- sulphide is used in the manufacture of sulphur dyes, and calcium poly sulphide, as a fungicide. EXERCISES 1. Calculate the percentage by weight of hydrogen sulphide contained in a solution of the gas in water saturated at 0. 2. What weight of ferrous sulphide which is 90 per cent pure is required to make enough hydrogen sulphide to fill a tank holding 5 cu. ft. with the gas at 4 atmospheres pressure at 0? 3. By what chemical reactions could you convert (a) zinc sulphide into zinc oxide, (6) copper oxide into copper sulphide, (c) ferrous sulphide into ferrous sulphate, (d) copper sulphide into copper, (e) lead into lead sul- phide? 4. Brass contains copper and zinc. How could the metals (a) be brought into solution, and (6) separated through the use of hydrogen sulphide? (c) How could the free metals be obtained from the sulphides? CHAPTER XX THE OXIDES AND ACIDS OF SULPHUR 285. On account of the fact that sulphur burns and occurs free in nature, sulphur dioxide has been known from the earliest times. It was used in certain religious ceremonies by the Greeks, and its disagreeable properties, no doubt, led to the selection of fire and brimstone in depicting the horrors of hell. Priestley first isolated the gas in pure condition as the result of an accident. He was attempting to distill sulphuric acid when his thermometer, which contained mercury, broke, and sulphur dioxide was formed. Priestley in his writings emphasizes the fact that many of his most important discoveries were made as the result of accident, but it should be noted that he had the ability to discern the important result of the accident and interpret it. Lavoisier explained the behavior of sulphur dioxide with water in 1777, and indicated the relation between the acid formed and sulphuric acid by naming the former sulphurous acid. 286. Occurrence of Sulphur Dioxide. This compound is found in the gases which issue from volcanoes. Great quantities of it are introduced into the air as the result of the burning of coal which contains sulphur. The gas which causes the disagreeable choking sensation in poorly ventilated railroad stations is sulphur dioxide. An important ore of copper is a sulphide; when this is smelted the sulphur is gotten rid of by heating it in air, and large quantities of sulphur dioxide are formed. As a result, in the neighborhood of smelters vegetation is destroyed for miles. Laws have been passed which prohibit the introduction into the air of the products of the furnaces which contain more than a definite amount of sulphur dioxide. In order to comply with these requirements it was neces- sary to find some way to dispose of the obnoxious gas. A profit- able solution of the problem was found in converting it into sul- phuric acid, which is a valuable product; in one of the largest 256 THE OXIDES AND ACIDS OF SULPHUR 257 sulphuric acid plants in the world the source of the sulphur is the waste gases obtained in the extraction of copper from a sulphide ore. 287. Preparation of Sulphur Dioxide. For industrial purposes sulphur dioxide is prepared by burning sulphur or heating a sul- phide in air. Iron pyrites, FeS 2 , is generally used for this purpose: 4FeS 2 + 11O 2 = 2Fe 2 O 3 + 8SO 2 Sulphur dioxide can be conveniently prepared by the action of an acid on a sulphite, which is a salt of sulphurous acid, H 2 S0 3 . When sodium sulphite and hydrochloric acid are used the reactions represented by the following equations take place : Na 2 SO 3 + 2HC1 = 2NaCl + H 2 SO 3 H 2 SO 3 =H 2 O The reaction resembles that between sodium carbonate, Na 2 COs, and the acid; in each case an acid is formed which breaks down into its anhydride and water. In the laboratory a saturated solu- tion of sodium bisulphite is commonly used, as it can be bought at a low price, and concentrated sulphuric acid is allowed to run into it slowly from a dropping funnel. The equation for the reac- tion is as follows : NaHSO 3 + H 2 SO 4 = NaHSO 4 + SO 2 + H 2 O Sulphur dioxide can also be prepared by reducing sulphuric acid. When the acid is heated with certain elements which have an affinity for oxygen, 1 atom of the latter is lost and the sulphuric acid, H 2 S04, is reduced to sulphurous acid, H 2 SO 3 , which in turn breaks down into sulphur dioxide and water. The reaction when copper is used can be considered to take place according to the steps represented by the following equations: H 2 SO 4 = H 2 SO 8 + O (1) . H 2 SO 3 = SO 2 + H 2 O (2) Cu + O = CuO (3) CuO + H 2 SO 4 = CuS0 4 + H 2 (4) 258 INORGANIC CHEMISTRY FOR COLLEGES The reactions thus indicated take place simultaneously. Oxygen is not set free as such, but unites with the copper to form copper oxide, which, being the oxide of a metal, reacts in turn with some of the sulphuric acid to form copper sulphate. It will be recalled that if oxygen gas is produced by a chemical reaction the formula used for it in the equation expressing the reaction is O2. In the first and third equations given above the oxygen is represented by the symbol O in order to indicate that these equations represent a process of oxidation and that none of the gas is set free. All processes of oxidation may be considered as taking place in two steps first, the breaking down of the oxidizing agent to furnish oxygen, and, second, the oxidation of the element or compound present. The four equations represent a single process and may be combined into one. The method by which this is done will be described fully as many reactions will be met with which are best considered as the result of the summation of several changes taking place at the same time. In a series of partial equations like those under discussion, if a symbol or formula appears on one side of one equation and on the other side of another, it indicates that the substance is first formed and then reacts. As a consequence, it is not present in the final product and its formula does not appear in the completed equa- tion for the reaction. In the case given above, oxygen for the reaction is furnished by the sulphuric acid; and it is used up by combining with the copper. We are justified, then, in canceling the O on the right-hand side of the first equation and the O on the left-hand side of the third equation. For the same reason the H2SO3 in equations 1 and 2, and the CuO in equations 3 and 4 can be canceled. The next step in arriving at the final equation is to add all the formulas that remain on the left side of the partial equations and make them equal the sum of the uncanceled for- mulas on the right side. The process yields the following, which is the equation that represents the reaction between copper and concentrated sulphuric acid: Cu + 2H 2 SO 4 = SO 2 + CuSO 4 + 2H 2 O In order to check the result, the final equation should be inspected to see if only the formulas of the substances used appear on one side and only the formulas of those actually obtained appear on the THE OXIDE AND ACIDS OF SULPHUR 259 other. The number of atoms of each element appearing on the two sides should also be noted to see if the equation balances. The process described above may appear to be a long and involved one to arrive at such a simple result. It is probable that the student could have written the final equation directly, if he had known what were the products formed in the reaction. The equations used as an example in this case are very simple ones, but others will be met with later which are more complicated and can be interpreted best by the use of the method described. 288. Sulphur dioxide is formed when concentrated sulphuric acid is heated with a number of the less active metals, for example, lead or mercury. In the case of the more active metals the acid is reduced farther. Copper removes but 1 oxygen atom, as we have seen, and H 2 SO3(SO 2 -f H 2 O) is formed. When zinc reacts with it, however, all the oxygen is removed and hydrogen sulphide is produced. Sulphuric acid does not oxidize the least active metals, such as platinum and gold. For this reason pans of platinum are used when dilute sulphuric acid is concentrated by boiling off the water which it contains. Sulphur dioxide is formed when a number of non-metallic elements are heated with concentrated sulphuric acid. With sulphur the reaction is expressed by the following equations: H 2 SO 4 = H 2 S0 3 + O H 2 SO 3 = SO 2 + H 2 O S + 2O = S0 2 There is not a fourth partial equation in this case, as there is when copper is used, because sulphur dioxide is not an oxide of a metal and does not, therefore, react with the sulphuric acid present to form a sulphate. Before the equations are combined the first one must be multiplied by 2, since 2 oxygen atoms are required for the oxidation of a sulphur atom, according to the third equation. Since no oxygen is set free there must be the same number of atoms of the element on the two sides of the equations. As the result of this change two molecules of sulphurous acid, H 2 SO3, are produced. This fact makes it necessary to multiply the second equation by 2, since all the sulphurous acid breaks down into sulphur dioxide and water. The equations can now be rewritten. It adds to the clear- 260 INORGANIC CHEMISTRY FOR COLLEGES ness to enclose in brackets the formulas of the substances which do not appear in the final equation. With these changes the equa- tions are as follows : 2H 2 SO 4 - [2H 2 S0 3 ] + [2O] [2H 2 SO 3 ] = 2SO 2 + 2H 2 O S + [2O] = SO 2 The addition of these partial equations leads to the following equation, which expresses the reaction between sulphur and con- centrated sulphuric acid: 2H 2 SO 4 + S ' = 3SO 2 + 2H 2 O 289. Physical Properties of Sulphur Dioxide. Sulphur dioxide is a colorless gas, which is slightly more than twice as heavy as air. It can be readily recognized by its odor, which is that produced when sulphur burns. The gas can be condensed to a colorless liquid by pressure alone, since its critical temperature is 156. The liquid freezes at 73 and boils at 10; it can be readily obtained from the gas by passing the latter through a vessel sur- rounded by a mixture of ice and salt. Liquid sulphur dioxide is put on the market in stout iron cans. This is possible because the pressure of the gas generated from the liquid at ordinary temperatures is not very great; at 20 it is 3.24 atmospheres. The gas is relatively soluble in water; 50 volumes under ordinary conditions dissolve in 1 volume of water. 290. Chemical Conduct of Sulphur Dioxide. The heat pro- duced when solid sulphur burns is 71,000 calories for each gram- atom (32 grams) of sulphur. As this number is relatively large we should expect sulphur dioxide, the product of the reaction, to be stable toward heat. It dissociates into its constituents only slightly at high temperatures, and in this behavior resembles water and other stable compounds. The chief chemical properties of sulphur dioxide are due to the fact that sulphur in many of its compounds has a valence of 6; in sulphur dioxide the valence of the element is 4, and, as a result, when the compound is brought into contact with certain other substances chemical reaction takes place. With chlorine, for example, sulphuryl chloride, SO 2 C1 2 , is formed by direct combination, the reaction being facilitated by THE OXIDES AND ACIDS OF SULPHUR 261 the presence of a trace of camphor, which serves as a catalytic agent. Sulphuryl chloride is a liquid which boils at 69; it decom- poses rapidly with water to form sulphuric and hydrochloric acids. Sulphur dioxide enters into a number of reactions of this type, the most important of which is the one that takes place between the gas and oxygen. As a result of the union sulphur trioxide, SOs, is formed. The reaction is one of great interest and will be discussed fully later, as it is the basis of one of the most important of the chemical industries the manufacture of sulphuric acid. Sulphur dioxide unites with water to form sulphurous acid: SO 2 + H 2 O < H 2 SO 3 The reaction indicated by the equation above is a reversible one; in an aqueous solution of sulphur dioxide there is always present some of the free oxide and some of the acid. It has been found impossible to isolate the latter on account of the fact that it decom- poses when an attempt is made to separate it from solution. The sulphurous acid present in the solution takes up oxygen from the air and from certain compounds that contain the element, and sulphuric acid, H 2 SO4, is formed. This reaction of sulphur dioxide, in which it is a reducing agent, is made use of in the process of bleaching, which will be described later. When brought into contact with certain substances sulphur dioxide loses its oxygen; with these it acts as an oxidizing agent. Only the most powerful reducing agents, however, can effect the reduction. One case has already been noted the reaction between sulphur dioxide and hydrogen sulphide : 2H 2 S -f SO 2 = 3S + 2H 2 O In this case the hydrogen sulphide is oxidized and the sulphur diox- ide reduced. Hydriodic acid is a powerful reducing agent which removes oxygen from sulphur dioxide: 4HI + SO 2 = 2H 2 O + S + 2I 2 291. Uses of Sulphur Dioxide. Large quantities of the gas are used in the manufacture of sulphuric acid, which is of fundamental importance in many chemical industries. 262 INORGANIC CHEMISTRY FOR COLLEGES Sulphur dioxide finds another important use in the bleaching of straw, wool, silk, and other substances which are affected dele- teriously by chlorine. The chemistry of the process has already been briefly described. In the presence of moisture the sulphurous acid first formed withdraws oxygen from the coloring matters present; as a result of the reduction colorless compounds are formed, and the product is said to be bleached. It is probable that in the case of straw and certain other substances, the bleaching is produced as the result of the union of sulphur dioxide with the colored compounds present. The resulting product is colorless but after exposure to light and the weather for some time, it undergoes decomposition and the original coloring matter is reformed. For this reason straw hats, which are bleached by exposure in a closed room to sulphur dioxide from burning sulphur, become yellow with age. When wool is bleached it is first scoured to remove adhering grease and then, while still moist, it is subjected to the action of sulphur dioxide produced by burning sulphur. A more convenient process, which is often used, is to pass the wool through a dilute solution of a sodium bisulphite and then through one of sulphuric acid. The acid decomposes the salt and liberates sulphur dioxide throughout the fiber. Sulphur dioxide was formerly much used as a disinfectant, as it kills the germs of disease. The gas was produced by burning sulphur in the room to be disinfected after windows and doors had been sealed. As the gas often bleached colored materials in the room, and more or less sulphuric acid was formed, the process was given up and formaldehyde was substituted for the purpose. Sulphur dioxide is used as a preservative, as it prevents the growth of bacteria. For this purpose it has been used in wines and beer and in preserving fruit which had to be shipped long dis- tances. Even small quantities of sulphur dioxide are said to be deleterious to health and the practice is now prohibited by law. Large quantities of sulphur dioxide are used in making salts of sulphurous acid. The most important of these are the acid sodium salt, NaHSOs, which is a source of sulphur dioxide for bleaching, as has been mentioned, and the corresponding calcium salt which is used in converting wood into pulp for the manufac- ture of paper. THE OXIDES AND ACIDS OF SULPHUR 263 292. Sulphurous Acid. It has already been stated that sul- phur dioxide when dissolved in water reacts with the solvent and is, in part, converted into sulphurous acid. The solution shows the characteristic properties of the solutions of all acids; it has a sour taste, turns blue litmus red, reacts with metals, and forms salts with bases. The chemical conduct of sulphur dioxide in this respect recalls that of carbon dioxide. Like the carbonates, the salts of sulphurous acid the sulphites are decomposed by other acids and the anhydride of the acid (SO 2 ) is liberated. Equa- tions for two typical reactions which illustrate the analogy are as follows : Na 2 C0 3 + 2HC1 = 2NaCl + CO 2 + H 2 O Na 2 S0 3 + 2HC1 = 2NaCl + SO 2 + H 2 O The explanation of the reaction in the case of sodium sulphite is similar to that offered for the formation of carbon dioxide from a carbonate (200); sulphurous acid is first formed as the result of double decomposition and then breaks down into sulphur dioxide and water. Sulphurous acid is a dibasic acid; the two hydrogen atoms present in it can be replaced by metallic atoms. The salts of the acid are called sulphites in accordance with the system of nomen- clature adopted by chemists. If the name of the acid ends in the syllable ous, the name of the salt ends in the syllable ite. Thus, sulphurous acid furnishes sulphites, chlorous acid chlorites, and arsenous acid arsenites. 293. Chemical Conduct of Sulphurous Acid. The solution of sulphur dioxide in water, which contains sulphurous acid, is an active reducing agent. The reaction which takes place between it and chlorine is typical: C1 2 + H 2 SO 3 + H 2 O = H 2 SO 4 + 2HC1 In this case the sulphurous acid has been oxidized to sulphuric acid, and the chlorine reduced to hydrochloric acid. The salts of sul- phurous acid behave in a similar way; they are used to destroy the excess of chlorine when the latter is employed as a bleaching agent. 294. Sulphites. There are two types of sulphites acid salts in which but 1 hydrogen atom of sulphurous acid, H 2 S03, has been 264 INORGANIC CHEMISTRY FOR COLLEGES replaced by a metallic atom, and neutral or normal salts in which both hydrogen atoms have been replaced. The method of naming these salts is the same as that used in the case of carbonic acid. The salt of the formula NaHSOs is called sodium hydrogen sul- phite, acid sodium sulphite, or sodium bisulphite. A number of sulphites are used commercially. Reference has been made to the sodium and calcium salts. The normal so- dium salt, Na2SOs, either anhydrous or in the hydrated form, Na2SO3,5H2O, is used in photography as a constituent of develop- ing solutions for films or paper. These solutions are powerful reducing agents and take up oxygen rapidly from the air. If this is permitted they soon lose their power as developers and turn dark and stain the film. If a sulphite is present this is largely prevented, because it absorbs the oxygen, being changed into a sulphate, and thus protects the developing agent from oxidation. Potassium metabisulphite is often used in photographic work instead of the neutral sulphite. It has the formula K2S2O.5, and may be considered as formed from 2 molecules of potassium bisulphite as the result of the loss of 1 molecule of water: 2KHS0 3 = K 2 S 2 5 + H 2 O The prefix meta is often used in inorganic chemistry in naming acids produced from other acids as the result of the loss of water; for example, HaPCU is the formula of phosphoric acid and HPOs that of metaphosphoric acid; the latter is formed as the result of the loss of a molecule of water from the former. Potassium metabisulphite crystallizes well, is relatively stable when dry, and weight for weight absorbs more oxygen than sodium sulphite. The solubility of the normal sulphites of the common metals resembles that of the carbonates. All are insoluble or difficultly soluble in water except those of sodium, potassium, and ammonium. The action of heat on most sulphites brings about a decomposi- tion into sulphur dioxide and the oxide of the metal present. The temperature at which this takes place varies with the metal. As is the case with the carbonates the sulphites of the less active metals are most readily decomposed. The salts of sodium and potassium decompose at very high temperatures only and in these cases a mix- ture of sulphide and sulphate is formed. If the heating is done in the presence of air, sulphites take on oxygen and are converted THE OXIDES AND ACIDS OF SULPHUR 265 into sulphates, which, in general decompose at a higher tempera- ture than sulphites. The test for sulphites is carried out by treating the compound with dilute hydrochloric acid and noting the odor of the evolved gas. SULPHUR TRIOXIDE 295. Preparation of Sulphur Trioxide. When sulphur burns in the air, the product is sulphur dioxide, but it is highly probable that more or less of the trioxide is also formed. At the high tem- perature of the flame the latter decomposes into the dioxide and oxygen. In confirmation of this view it has been shown that if the air in which the sulphur burns contains water-vapor, appreciable quantities of the trioxide can be detected as the result of the reaction. The amount of sulphur trioxide formed when sulphur burns is small, and the reaction does not serve, therefore, as a means of preparing it. Sulphur trioxide is formed when the sulphates of certain metals are heated. The substance was formerly prepared in this way, but it is now produced directly as the result of the union of sulphur dioxide and oxygen of the air in the presence of platinum. As the preparation of the oxide in this way is the most important step in one of the methods used to manufacture sul- phuric acid, it will be described in detail later. 296. Physical Properties of Sulphur Trioxide. There are two forms of sulphur trioxide; the one obtained as the result of the union of sulphur dioxide and oxygen in the presence of finely divided platinum and in the absence of water, is a colorless liquid which boils at 46 and freezes at 15 to a solid resembling glass in appearance. If a trace of moisture is added, some sulphuric acid is formed, and, as a result of the presence of a small amount of acid, the liquid changes on standing to a mass of white, opaque, needle- shaped crystals. These do not melt when heated, but pass into a vapor, which on condensation yields the liquid variety of the tri- oxide. The molecular weights of the two forms have been deter- mined by the freezing-point method; from these it has been cal- culated that the formula of the liquid variety is SOs and that of the solid form 82(^6. The white crystalline variety of sulphur trioxide is said to be a polymer of the liquid form. It should be 266 INORGANIC CHEMISTRY FOR COLLEGES called sulphur hexoxide, as it is really a different substance from sulphur trioxide. Since it passes so readily into the trioxide and its chemical behavior is identical with that of the latter, the dif- ference is not emphasized in its name. 297. Chemical Behavior of Sulphur Trioxide. The combining power of an element, as indicated by its valence, is one of its most important chemical characteristics. It determines the composi- tion of the compounds derived from the element, and, in certain cases, even its fundamental chemical behavior; some elements have acid-forming properties when they exhibit one valence, and base-forming properties when they show another. The processes of oxidation and reduction are associated with change in valence, and many other important facts of chemical interest can be inter- preted through a study of the valencies of the elements involved. A comparison of the behavior of sulphur dioxide and sulphur trioxide when heated to high temperatures brings out a conclusion which can be applied to other compounds. The dioxide is very stable towards heat; the trioxide, on the other hand, begins to decompose below 400 into the dioxide and oxygen, and as the tem- perature is raised the decomposition rapidly increases. The reaction which takes place is a reversible one, 2SO 3 ^ 2SO 2 + O 2 and the equilibrium is markedly affected by the temperature. The facts in regard to this reaction have been carefully studied, since it is the basis for the manufacture of sulphuric acid, as has been repeatedly mentioned. The particular fact to be emphasized at this point is that the combining power of an ele- ment is markedly affected by temperature; with rise in tempera- ture the ability of an element to hold others in combination with it, falls off. This is a generalization which applies to all elements, and is, therefore, of prime importance. Sulphur trioxide reacts with water violently; when a small amount of it is dropped into the liquid a hissing sound results, like that produced when red-hot iron is thrust into water. Since sulphuric acid is the product of the reaction sulphur trioxide is often called sulphuric anhydride: SO 3 + H 2 = H 2 S0 4 THE OXIDES AND ACIDS OF SULPHUR 267 On account of its great affinity for water sulphur trioxide fumes in the air (141). When the vapor comes into contact with moist air a cloud of great density is formed. This fact was utilized in the recent war to produce smoke screens for use along the battle front and to protect vessels at sea. To produce a dense cloud on land sulphuric acid containing sulphur trioxide in solution, so called fuming sulphuric acid, was allowed to run on lime. The reaction which took place between the acid and the base produced sufficient heat to volatilize the sulphur trioxide, which produced a dense white smoke with the moisture in the air. To make a smoke screen at sea the fuming acid was introduced into the smoke-stack of the vessel where the hot gases from the fires under the boiler vaporized the sulphur trioxide. SULPHURIC ACID 298. Sulphuric acid is a fundamental raw-material in many of the largest chemical industries, and in others in which it is not itself used, substances are employed that require it in their prep- aration. Sulphuric acid is thus the most widely used manufac- tured compound in industrial chemistry. Liebig, one of the leaders in the early days of chemistry, said the civilization of a nation could be gauged by the amount of soap it consumed per capita; later it was claimed that sulphuric acid should be the standard of measure, for we can trace back to it so many of the essentials and conveniences of material existence. The manufacture of the acid has been one of the great problems of industrial chemistry, and many chemists have exercised their skill and ingenuity in the solu- tion of the problem, which is a very complicated one. Sulphuric acid was known in ancient times because it is pro- duced by the action of water on the sulphur trioxide formed when many naturally occurring sulphates are heated to a high tempera- ture. Fuming sulphuric acid (H 2 SO4 + SOs) was manufactured in the middle of the eighteenth century at Nordhausen, in Ger- many, by heating ferrous sulphate, FeSO4, 7H2O, but the product was very expensive. The common name of the salt was iron vitriol, and the oily product produced from it was called oil of vitriol a name still used in chemical trade. In 1740 Ward, in England, started the manufacture of the acid by burning sulphur 268 INORGANIC CHEMISTRY FOR COLLEGES with niter, KNOs, in the presence of water under great glass domes. A few years later lead chambers were substituted for the glass domes, and in 1793 it was shown that the oxidation of sulphur dioxide in the process was brought about through the catalytic influence of the oxides of nitrogen formed from the nitrate used. 299. Manufacture of Sulphuric Acid. The preparation of the acid involves, in the main, but a few simple reactions; sulphur, or a sulphide, is burned to sulphur dioxide, which is then oxidized, and the resulting compound is treated with water. It has already been shown how this can be done, but serious difficulties are met with when the process is carried out on an industrial scale. The reac- tion as the result of which sulphur dioxide unites with oxygen and forms sulphur trioxide in the presence of catalytic agents, has been known for a long time, but it was only in 1901 that it was applied to the manufacture of sulphuric acid. As the result of extended chemical research a method was worked out for the synthesis of indigo our most important blue dye. This method involved, in one of the steps, the use of fuming sulphuric acid. To compete successfully with the natural dyestuff obtained from a plant grown in India, it was necessary to reduce the cost of the production of sulphur trioxide. In a search for a new method, the direct oxida- tion of sulphur dioxide to the trioxide in the presence of catalytic agents was carefully studied, and as the result the so-called " contact " process for the manufacture of sulphuric acid was worked out. While the method was primarily devised for the preparation of sulphur trioxide and fuming sulphuric acid, it was readily applied to the manufacture of the ordinary acid. The older, more cumbersome process is still in use to-day, but it is probable that as new plants are erected the contact process will be employed, except in the case where the nature of the raw materials available favor the use of the older method. The oxidation of sulphur dioxide to trioxide takes place very slowly under ordinary conditions and it is necessary to increase the rate of the reaction when it is used in a manufacturing opera- tion. In the contact process the catalytic agent which increases the rate is platinum; in the older, so-called " chamber " process certain oxides of nitrogen are used for this purpose. The reactions which take place in the latter process are more or less complicated, as we shall see, but in the main they involve the oxidation of the THE OXIDES AND ACIDS OF SULPHUR 269 sulphur dioxide by nitrogen dioxide, NC>2, which is reduced, as a result, to nitric oxide, NO. This oxide unites directly with the oxygen in the air to form nitrogen dioxide, which, in turn, oxidizes more sulphur dioxide. The nitric oxide serves thus as a carrier of the oxygen from the air to the sulphur dioxide; at one instant the nitrogen is in the form of NO and at the next of NO2, then NO, then NO2, and so on. The word carrier has been selected to name this process, which, it will be seen, is a catalytic one, for the nitrogen compounds involved do not enter into the final product of the reaction. 300. Contact Process for Sulphuric Acid. The equation for the reaction upon which this process is based is as follows : 2SO 2 + O 2 =* 2SO 3 + 41,800 cal. The equation shows that the reaction is reversible, and indicates the amount of heat evolved when it takes place. Both of these facts are of prime importance. To get the best result, the reac- tion must be carried out under the conditions where the equilib- rium is such that the largest practical amount of sulphur trioxide is formed. These conditions are determined by the temperature at which the reaction takes place; and since heat is evolved, the proper conditions can be obtained only by carrying the process out in such a way that the heat can be controlled. Before the reaction can be adapted to the commercial preparation of sulphur trioxide, the effect of temperature on the equilibrium must be known, and the effect of various catalyzing agents on the rate at which the reaction proceeds must be determined. At the tempera- ture at which the reaction proceeds rapidly enough in the absence of a catalyzer to make it useful, the equilibrium is such that but a small amount of sulphur trioxide is present in the gases. At 400 when twice the theoretical amount of oxygen is used, between 98 and 99 per cent of the dioxide is converted into trioxide; at 720, the conversion is about 60 per cent complete; and at 900 prac- tically no sulphur trioxide is formed. It is evident that the reac- tion should be carried out at approximately 400 if possible. A large amount of heat is produced as a result of the reaction, and as the temperature would rise above 400 if the heat were not controlled, it is necessary to construct the apparatus in which the reaction is carried out in such a way that the excess heat over that 270 INORGANIC CHEMISTRY FOR COLLEGES required to maintain the proper temperature is carried away. This is done by using this excess to pre-heat the gases to the tem- perature at which the reaction takes place. A number of substances were studied in the investigation of the influence of catalytic agents on the oxidation of sulphur dioxide, and it was found that finely divided platinum and an oxide of iron, Fe2Oa, were the only ones that were of practical significance. Catalyzers have no effect on the final equilibrium attained in a reaction, but differ from one another in their influence on the rate of the reaction. The oxidation of sulphur dioxide takes place rapidly at 400 in the presence of platinum, and at 625 when ferric oxide, Fe 2 C>3, is used. Since at the latter temperature only about 70 per cent of the dioxide is converted into trioxide, it is evident that platinum is preferred as a catalytic agent notwith- standing its high cost. In some plants two contact chambers are used; the first contains ferric oxide and the second platinum. 301. A diagrammatic sketch of the lay-out of a plant for the manufacture of sulphuric acid by the contact process, a so-called WashingTpwer ContactTower Absorption Tower I DryingTower j \H 2 SO 4 Sulfur Burner / Conc } (0=$. I Dust Chamber / ... . H 2 S0 4 / , Jfc Air FIG. 29. flow-sheet, is given in Fig. 29. The various steps in the process can be seen by an examination of the diagram. The gases issuing from the contact chamber are passed through strong sulphuric acid (about 97 per cent), which absorbs the sul- phur trioxide. 302. Lead Chamber Process for Sulphuric Acid. This process, which has been the standard one for over a hundred years in chem- ical industry, derives its name from the fact that the sulphuric THE OXIDES AND ACIDS OF SULPHUK 271 acid is formed in chambers built of lead. The mixing of the gases involved takes place rather slowly in the presence of the large amount of nitrogen in the product resulting from the burning of sulphur in the air; and, as a consequence, the reaction chambers must be large. They are built of lead, as this metal resists satisfactorily the acid formed in the process. Fig. 30 shows diagrammatically the principal features of a sulphuric acid plant. Sulphur dioxide is first produced by burning * < * -NOandNO z TowerAcid. Nffmsyl 5u Acid s ^JaoQQ~r\ I L-fstLff*** ^ft^S^ tjyr c ' er r oxn FIG. 30. sulphur or by roasting a sulphide, usually pyrite, FeS2, in the pres- ence of air: 4FeS 2 + 11O 2 = 2Fe 2 O 3 + 8SO 2 The gases from the burners traverse a flue in order to allow the dust present to settlej and are there mixed with the required amount of air, which furnishes the oxygen for the subsequent oxidation. In the flue is placed a pot containing sodium nitrate and sulphuric acid, which furnish to the gases enough nitric acid to produce the small amount of the oxides of nitrogen equivalent to that lost in the subsequent oxidation. The gases next pass through the Glover tower, which will be explained later, pick up the oxides of nitrogen there, and then go into the lead chambers, where they come into contact with water, which is introduced at several places in the form of steam or as a fine spray. As the gases slowly drift through 272 INORGANIC CHEMISTRY FOR COLLEGES they are converted into sulphuric acid which collects as a liquid on the floor of the chambers. The large amount of free nitrogen present in the gases carries along with it the oxides of nitrogen. In order to prevent the loss of the latter the gas issuing from the chambers is passed through the Gay-Lussac tower; this contains tiles or coke over which a stream of concentrated sulphuric acid runs continuously. The acid absorbs the oxides of nitrogen, and on reaching the bottom of the tower is pumped to the top of the Glover tower. It is allowed to flow down through this tower, which is filled with pieces of broken flint, along with some of the dilute acid obtained from the lead chambers. The compound formed from the oxides of nitrogen and sulphuric acid in the Gay-Lussac tower is decomposed by the water present in the dilute acid, and the oxides thus set free. The hot gases from the pyrite burners in passing through the Glover tower carry along with them the oxides of nitrogen into the lead chambers. As the dilute acid falls through the Glover tower the hot gases in passing through it carry off most of the water present, and, as a result, the sulphuric acid which flows from the bottom of the tower contains but a small amount of water. This acid is pumped to the top of the Gay-Lussac tower and serves to collect the oxides of nitrogen as before. The final product is the acid which collects in the chambers the so-called chamber acid. 303. The chemical reactions which take place in the chamber process for the manufacture of sulphuric acid have been very fully studied. Several explanations of the way in which the oxides of nitrogen act have been put forward. While the part played by the oxides in effecting the oxidation of sulphur dioxide is essentially that already indicated, the reactions are more complex than those given. When an insufficient amount of water is present in the lead chambers during the oxidation, a crystalline compound is formed on the walls of the chambers. This substance is supposed to be an inter- mediate product in the reaction; it has the composition represented by the formula H(NO)SO4 and is called nitrosyl sulphuric acid. The formula is written as it is to show the relation between the composition of the substance and that of sulphuric acid, H 2 SO 4 ; it may he considered as formed by the replacement of one hydrogen atom in the acid by the NO group. As this group appears in a number of compounds it has been given a special name nitrosyl. The formulas of sulphuric acid and nitrosyl sulphuric acid are often written in the following manner: O=Q- O - H O=Q- O - H 0=- - H 0=- O - THE OXIDES AND ACIDS OF SULPHUR 273 It Is seen that the NO group replaces 1 hydrogen atom. Nitrosyl sulphuric acid can be prepared in a number of ways; for example, it is produced as the result of the direct addition of nitric acid and sulphur dioxide: SO 2 -f HNO 3 = H(NO)SO 4 It is formed when nitric oxide and nitrogen dioxide are passed into concen- trated sulphuric acid, (1) 2H 2 S0 4 + NO+ NO, < 2H(NO)SO 4 + H 2 O and when sulphur dioxide is oxidized by nitrogen dioxide in the presence of water: (2) H 2 O + 2SO 2 + 3NO 2 = 2H(NO)SO 4 + NO With these equations the chief reactions which take place in the formation of sulphuric acid can be interpreted. In the chamber, reaction (2) probably first takes place. Reaction (1) is reversible and, as a result, in the presence of water the nitrosyl sulphuric acid formed breaks down into sulphuric acid and the oxides of nitrogen. The air present converts the nitric oxide, NO, into the dioxide which serves to react with more sulphur dioxide. It was stated that nitrosyl sulphuric acid the so-called chamber crystals sepa- rates only when there is a deficiency of water in the chamber; Equation (1) furnishes a reason for this. The gases swept along by the nitrogen pass through the Gay-Lussac tower. Here the reaction represented by Equation (1), reading from left to right, takes place. It is seen from this why concentrated sulphuric acid is needed at this stage. The product from this tower, which contains nitrosyl sulphuric acid in solution, is sent through the Glover tower, where it comes in contact with diluted sulphuric acid. The water present decomposes the nitrosyl sulphuric acid, as indicated by Equation (1), reading from right to left; and the oxides of nitrogen liberated are carried along with the sulphur dioxide and oxygen into the contact chambers, where the cycle is begun once more. By considering Equations (1) and (2), along with the equation for the oxidation of nitric oxide to nitrogen dioxide, as partial equations in a single transformation, we can bring out clearly the fact that the nitrogen compounds are catalytic agents. The three equations written in this way, the formulas of the compounds appearing on both sides of the equation being enclosed in brackets as usual, are as follows: H 2 0+ 2S0 2 + [3N0 2 ] - [2H(NO)S0 4 ] + [NO] [2H(NO)SO 4 ] + H-O= 2H 2 SO 4 + [NO]-f- [NOJ [2NO] + O 2 = [2NO 2 ] 2H 2 O + 2SO 2 + O 2 = 2H 2 SO 4 274 INORGANIC CHEMISTRY FOR COLLEGES The equations differ from the partial equations studied before because, in this case, they represent reactions that can be independently realized; the oxidation of sulphur dioxide represented by the combined equation can be carried out in the separate steps indicated by the partial equations. These partial equations show clearly how nitrogen dioxide serves as a catalytic agent in effecting the oxidation. In the case of the use of platinum a satisfactory explanation is not so evident; various views have been put forward, but adequate experimental proof for any one is lacking. It was thought at one time that the platinum formed an unstable compound with oxygen, and that this oxidized the sulphur dioxide and was, as a result, reduced to platinum; the metal united with more oxygen and the cycle was repeated. According to this explanation the platinum serves as a carrier in the same general way that nitric oxide does. A later explanation is the more probable one. According to this, the finely divided metal adsorbs large amounts of the gases on its surface and under these conditions reaction between them takes place more rapidly. This view serves to explain the behavior as catalytic agents of very inert substances like powdered glass and sand, where the formation of an unstable chemical compound with the reacting substances is highly improbable. 304. The product of the chamber process the so-called chamber acid contains about 65 per cent of sulphuric acid and varies in specific gravity between 1.5 and 1.6. In this form the acid is Used in a number of commercial operations, but for others, it is concentrated by evaporating off most of the water. This can be carried out in vessels constructed of lead until the con- centration reaches 77 per cent, when the specific gravity is 1.7. The acid reacts slowly with lead below this temperature, but the lead sulphate formed adheres closely to the metal and serves as a protective coating. Hot sulphuric acid stronger than 77 per cent dissolves lead sulphate, and, as a consequence, the acid cannot be concentrated beyond this point in vessels made of lead. The final evaporation is carried out either in silica pans or stills made of cast iron; when the concentration reaches about 94 per cent and the acid has the specific gravity 1.84 the evaporation is stopped. The product is the concentrated sulphuric acid of commerce. Sulphuric acid is shipped in steel drums or tank cars, or in carboys, which are large glass bottles packed in straw in wooden boxes, the neck of the bottle being exposed so that the acid can be conveniently poured out. The usual form of carboy holds about 200 pounds of sulphuric acid. 305. Physical Properties of Sulphuric Acid. The pure, anhy- drous acid, sometimes called hydrogen sulphate, is an oily liquid THE OXIDES AND ACIDS OF SULPHUR 275 which has the specific gravity 1.84 at 15, is miscible with water in all proportions, and freezes to a crystalline solid that melts at 10.5. When heated to its boiling-point, 336, it undergoes partial decomposition and a mixture of sulphuric acid and water containing 98.33 per cent of the acid is obtained. This mixture boils at 338 and is the product obtained when the acid is distilled. Commercial sulphuric acid contains a number of impurities; among these are lead sulphate, arsenic oxide, and oxides of nitrogen. The presence of lead sulphate can be shown by diluting the acid with water; it is precipitated because it is insoluble in dilute sulphuric acid. The so-called C.P. (chemically pure) acid con- tains but very small amounts of impurities, the presence of which can be neglected in most of the cases in which the acid is used. 306. Chemical Behavior of Sulphuric Acid. Sulphuric acid is an active chemical reagent. It shows a number of reactions of addition. It unites directly with water with the evolution of a large amount of heat. For this reason care must be taken when the acid is diluted with water. If the latter is poured on the acid it floats on the surface and the large amount of heat generated causes the formation of steam so rapidly that a lively sputtering results; if, on the other hand, the acid is poured into the water, it sinks, the two liquids mix rapidly, and there is no chance of an accident. Several hydrates of sulphuric acid have been isolated; the one having the formula H^SO^B^O is a crystalline compound which melts at 8; at very low temperatures hydrates having the composition represented by the formulas H2SO4,2H20 and H2SO4,3H2O are formed. The affinity of sulphuric acid for water is so great that it is an excellent agent for removing water-vapor from gases that do not react with it. To effect this the gas is simply bubbled through the liquid. Sulphuric acid decomposes certain compounds containing hydrogen and oxygen and withdraws from them these elements in the proportion to form water. When the concentrated acid, for example, is warmed with sugar, C^H^On, the principal reaction which takes place is the liberation of carbon as a black mass, and the formation of water which unites with the sulphuric acid. A similar reaction takes place when the acid is allowed to stay in contact with wood. 276 INORGANIC CHEMISTRY FOR COLLEGES Sulphuric acid reacts with sulphur trioxide and forms disul- phuric acid, H 2 SO 4 + SO 3 <= H 2 S 2 O 7 which is present in the so-called fuming sulphuric acid; it readily breaks down into its constituents and the sulphur trioxide fumes when it comes in contact with moist air. Salts of the acid, how- ever, are stable; they are usually called pyrosulphates 1 and are formed by heating acid sulphates to a high temperature : 2NaHSO 4 + Na 2 S 2 O 7 + H 2 O Sulphuric acid unites directly with normal salts of the acid and forms acid salts : H 2 SO 4 + Na 2 SO 4 - 2NaHSO 4 Sulphuric acid enters into reactions of double decomposition with salts of other acids, and is a very valuable reagent in the preparation of acids having boiling-points lower than that of sulphuric acid itself; since this is relatively high, many acids can be prepared in this way. The preparation of hydrochloric acid is an example that is already familiar. Sulphuric acid enters into reactions of oxidation. Some of these have been discussed in the consideration of the preparation of sulphur dioxide (287). All the metals except the so-called noble metals, among which are platinum and gold, can be oxidized by sulphuric acid. In all cases sulphates are formed as the final result of the action since if oxides were first produced they would react with the acid present to form salts. The product formed as the result of the reduction of sulphuric acid varies with the activity of the metal with which it interacts. The more active metals like zinc and magnesium reduce the acid to hydrogen sulphide: H 2 S0 4 = H 2 S + [4O] 4Zn + [4O] = [4ZnO] [4ZnO] + 4H 2 SO 4 = 4ZnSO 4 + 4H 2 O 4Zn + 5H 2 SO 4 = H 2 S + 4ZnSO 4 + 4H 2 O 1 The prefix pyro is derived from the Greek word signifying fire. It has been used often in naming compounds produced as the result of the action of heat. THE OXIDES AND ACIDS OF SULPHUR 277 The reaction which takes place, however, is more complicated than the equation indicates, for some of the hydrogen sulphide formed is oxidized by the sulphuric acid and sulphur is formed. Hydrogen reduces the acid to sulphur dioxide at about 160. Sulphuric acid oxidizes many of the non-metals. The reaction with sulphur has been given (288). When the acid is heated with carbon, carbon dioxide is formed: 2H 2 SO 4 = 2H 2 O + 2SO 2 + [2O] C + [20] - C0 2 2H 2 SO 4 + C = 2H 2 O + 2SO 2 + CO 2 The activity of sulphuric acid as an oxidizing agent increases with rise in temperature, and decreases as the acid is diluted with water. When sulphuric acid is dissolved in water it undergoes ioniza- tion : , H 2 S0 4 *=* H + + HSO 4 ~ HS0 4 - +* H + + S0 4 - The extent to which the acid is ionized is determined by the con- centration that is, by the relative amount of water present. In a one-tenth normal solution (4.9 grams in a liter) at 18, about 61 per cent of the acid is converted into hydrogen and sulphate, S0 4 ~~, ions. Salts of sulphuric acid form ions in a similar way; copper sulphate yields, for example, a copper ion, Cu ++ , and a sulphate ion, S0 4 ~~. 307. Sulphates. Normal sulphates of most of the metals are known, and many of them have important industrial applications, which will be considered later in connection with a study of the metals that they contain. The sulphates which occur in nature as minerals have a commercial significance and are mined in large quantities. The insoluble sulphates can be readily formed by double decomposition from aqueous solutions of salts, one of which contains the sulphate ion and the other the ion of a metal. The soluble sulphates are prepared by the action of sulphuric acid on oxides, hydroxides, or salts of acids which are more volatile than sulphuric acid, for example, carbonates and chlorides. 278 INORGANIC CHEMISTRY FOR COLLEGES The important acid sulphates are those of the so-called alkali metals, sodium and potassium. They can be formed by the par- tial neutralization of sulphuric acid; sodium hydrogen sulphate can be made in this way: NaOH + H 2 SO 4 = NaHSO 4 + H 2 O Acid salts are also formed by treating the normal salts with sul- phuric acid: Na 2 SO 4 + H 2 SO fi = 2NaHS0 4 These salts show an acid reaction in solution, since they yield hydrogen ions on dissociation. The sulphates of the heavy metals decompose at high tempera- tures into sulphur trioxide and the oxide of the metal present, for example, ZnSO 4 = ZnO + SO 3 The sulphates of sodium and potassium do not undergo a similar decomposition. Acid sulphates on heating lose water and pass into pyrosulphates: 2NaHSO 4 = Na 2 S 2 O 7 + H 2 O 308. Test for Sulphates. This test is based upon the fact that barium sulphate is the only common barium salt that is insoluble in dilute nitric acid. When a solution of a barium salt is added to a solution of sulphuric acid or a sulphate ; a white precipitate of barium sulphate a very insoluble substance is formed as a result of double decomposition : Na 2 SO 4 + BaCl 2 = 2NaCl + BaSO 4 When this equation is written in the ionic form, as follows, 2Na + + SO 4 -- + Ba ++ + 2C1" = 2Na + + BaSO 4 + 2C1~ it is seen that the reaction consists in the union of the barium and sulphate ions to form insoluble barium sulphate. The sodium chloride formed at the same time, being soluble, remains in solu- tion and is, consequently, in the form of ions. The test is, there- fore, based on the fact that if a barium ion is brought into contact with a sulphate ion, the two unite and form barium sulphate, because this salt is insoluble. There are many barium salts whi4~, are discharged by the current, the sodium liberated reacts with the water present and forms hydrogen and sodium hydroxide; the HSCU" ions liberated unite and form persulphuric acid, H 2 S 2 Og, which interacts by double decomposition with some of the sodium hydrogen sulphate to form sodium persulphate; and this being difficultly soluble separates out. The union of two discharged HSC>4~ ions to form H 2 S 2 Og is entirely analogous to the formation of chlorine molecules, C1 2 , THE OXIDES AND ACIDS OF SULPHUR 281 when Cl~ ions are discharged. A solution of persulphuric acid can be obtained by electrolyzing an aqueous solution of sulphuric acid containing about 50 per cent of the latter. It has been shown that the ammonium salt can advantageously replace the sodium salt in the preparation of persulphates, and large quantities are now manufactured in this way. Persulphuric acid is not known in the pure condition; when an attempt is made to prepare it from its salts it decomposes with the formation of either oxygen or hydrogen peroxide, depending on the conditions. This fact leads to the view that the acid is related to hydrogen peroxide, and that in it some of the oxygen atoms are joined as they are in the peroxide. It will be recalled that certain acids are formed as the result of the union of oxides with water; sulphur trioxide and water, for example, give sulphuric acid. Such acids can be broken down into oxides (anhydrides) and water. It is highly probable that hydrogen peroxide reacts with anhydrides in a similar way, for acids are known, the formulas of which bear the same relation to hydrogen peroxide that the formulas of the ordinary acids bear to water. For example, if sulphur trioxide re- acts with hydrogen peroxide as indicated by the following equation, SO 3 + H 2 O 2 = H 2 SO 5 we would expect the resulting acid to yield hydrogen peroxide as the result of decomposition. An acid of the above formula is known, called Caro's acid from its discoverer, which is prepared by the action of hydrogen peroxide on strong sulphuric acid. It is a powerful oxidizing agent. If 2 molecules of sulphur trioxide reacted with 1 of hydrogen peroxide the resulting acid would have the formula H 2 S 2 Os, which is that of persulphuric acid. The relationship indicated is probably true for the decomposition of persulphuric acid in water solutions yields either hydrogen peroxide or oxygen, which is probably formed from the peroxide: H 2 S 2 O 8 + 2H 2 O = 2H 2 SO 4 + H 2 O 2 This reaction is utilized in one of the methods used industrially for the preparation of hydrogen peroxide. Ammonium persulphate is treated with strong sulphuric acid; the persulphuric acid liberated breaks down spontaneously into hydrogen peroxide, which is '282 INORGANIC CHEMISTRY FOR COLLEGES obtained from the mixture by distillation under diminished pres- sure. Persulphuric acid and its salts are oxidizing agents, a fact which leads to their use in bleaching. They decompose slowly in solution, but rapidly in the presence of a substance which they can oxidize. In this decomposition sulphates, sulphuric acid, and oxygen are formed. Persulphates are used in " reducing " photographic negatives. When a negative has been over developed too much silver has been formed and the plate is too dense. As much silver can be removed from the negative as desired by putting it into a solution of a per- sulphate. The latter slowly oxidizes the metal and converts it into silver sulphate, which is dissolved by the water present: 2Ag + K 2 S 2 O 8 = K 2 S0 4 + Ag 2 S0 4 EXERCISES 1. Write equations for the reactions which take place between the follow- ing substances: (a) SO 2 C1 2 + H 2 O, (6) concentrated H 2 SO 4 and Ag, (c) CaSO 3 + H 2 SO 4 , (d) Na 2 S 2 O 7 + H 2 O, (e) concentrated H 2 SO 4 and Hg, (/) concentrated H 2 SO 4 and As as the result of which As 2 O 3 is formed. 2. Write the graphic formula of the compound formed as the result of the addition of HC1 to SO 3 , and an equation for the reaction of the product with H 2 O. 3. State two ways in which you could separate SO 2 from CO 2 . 4. Sulphurous acid is a much stronger acid than carbonic acid. What would happen if sulphuric acid were added to a solution containing sodium sulphite and sodium carbonate, the amount of the acid being less than that required to decompose either salt? 5. What would happen if dilute hydrochloric acid were added to a solu- tion of (a) sodium sulphide, (6) sodium thiosulphate, and (c) a mixture of the two salts. Write equations for all reactions. 6. What volume of oxygen at and 760 mm. will react with a solution containing 100 grams of (a) hydrated sodium sulphite, (6) the anhydrous salt, and (c) sodium metabisulphite? 7. When an excess of iodine is added to a dilute solution of sulphur dioxide the latter is oxidized quantitatively to sulphuric acid. Devise a volumetric method for the quantitative determination of SO 2 based on this reaction. 8. What weight of sulphur must be burned to produce enough SO 2 to saturate 1 liter of water with the gi:s at room temperature? 9. Starting with Na 2 SO 4 ,10H 2 O write equations for reactions by which the following could be prepared: (a) NaHSO 4 , (6) Na 2 S 2 O 7 , (c) Na 2 S, (d) Na 2 S0 3 , (e) Na 2 S 2 O 3 . 10. Write an equation for the reaction between Na2SOs and Na 2 S 2 O 8 . THE OXIDES AND ACIDS OF SULPHUR 283 11. A mixture of sodium sulphate and salt was analyzed with the following result: 1.000 gram of the mixture when dissolved in water and treated with barium chloride gave 1.315 grams barium sulphate. Calculate the percent- age of sodium sulphate in the mixture. 12. (a) How many cubic feet of air are necessary to convert 1 ton of pyrite, FeS 2 , into SO 2 ? (6) What is the relation between the volume of the SO 2 and N 2 if no excess of air is used? (1 pound molecular weight occupies 359 cu. ft.) CHAPTER XXI NITROGEN AND THE ATMOSPHERE 312. The air has aroused in man the greatest interest since the earliest days. It was feared and venerated, and like other things in nature which affected man's existence for good or evil it was deified. In early Grecian days sailors prayed and burned incense to the God of the Storm before setting out on a voyage. As men began to question nature they appreciated the significance of the air and it became one of the fundamental concepts in ancient philosophy. Empedocles, a Greek, put forward in the fifth cen- tury before Christ the view that everything was made of four elements, earth, air, fire, and water. This view persisted through centuries, but was modified as time went on and the air was more and more studied. John Mayow (1654-1679) showed that when certain metals were heated in the air they withdrew something from it; Priestley demonstrated in 1774 that the air contains something which is essential in combustion, and obtained this important substance in pure condition by heating the oxide of mercury; and, finally, Lavoisier explained the part played in combustion by Priestley's gas, which he called oxygen, and demon- strated by convincing experiments that his view was correct. It can be said with assurance that the science of chemistry was born when men first began to study experimentally the nature of the air. In the early history of the science of physics, also, the study of air plays an important part. The explanations first offered of physical phenomena were, in general, like those put forward in the case of chemical changes. The forces of nature were assigned human attributes; water could not be pumped up into a tube higher than 34 feet, because nature " abhorred " a vacuum. Tor- ricelli came to the conclusion that the rise of the water was due to the pressure of the air. If this were true the height of a col- umn of a heavy liquid supported by air would be less than that of 284 NITROGEN AND THE ATMOSPHERE 285 a column of water. Torricelli tried the experiment with mercury in 1643 and, as a result, invented the barometer. Pascal observed later that the height of a barometer on the top of a mountain was less than at its base, a fact that confirmed the view of Torricelli, for at the top there is less air to support the mercury. Guericke invented the air-pump in 1650; Boyle at once made one for him- self and as the result of his experiments with it discovered the law for which his name is famous. 313. The atmosphere has been one of the chief causes in shaping the earth's surface as it is to-day. The carbon dioxide and water- vapor of the air decompose rocks, slowly converting them into the materials of which the soil is composed. For this and other reasons the influence of the air and its constituents is studied in Geology. Meteorology the science of climate and the weather is primarily concerned with air from the physical point of view. The micro-organisms in the air bring about many chemical changes through the agency of fermentation; the putrefaction of animal substances, and the decay of vegetable material, such as the rotting of wood, are examples of important changes that are brought about in this way. Pasteur was able to prove the impossibility of spontaneous generation the production of life without the presence of living organisms by carrying out experiments in the absence of ordinary air. When the latter was freed from all bacteria, molds, etc., and when the medium in which the experiment was carried out was sterilized, no living thing was produced. NITROGEN 314. A long time before nitrogen was isolated, it had been observed that air contains two substances, one of which was inac- tive and did not take part in combustion. Rutherford in Scotland in 1772 recognized the gas as a distinct substance, and at about the same time Scheele in Sweden showed how it could be obtained by burning substances in the air in order to remove the oxygen; the residue after removal of the products of combustion was the substance we now call nitrogen. It was shown later that the element was present in niter (potassium nitrate, KNOs), and this fact led to the selection of the English name for the gas. Lavoisier proved that nitrogen is an elementary substance. 286 INORGANIC CHEMISTRY FOR COLLEGES 315. Occurrence of Nitrogen. About four-fifths of the air is free nitrogen. We have, thus, an unlimited supply of this element which plays such an important part in life-processes. Up to recent years, however, man was unable to utilize the air as a source of nitrogen compounds to be used in fertilizing the soil and for other important purposes, on account of the great chemical inactivity of the element. In certain natural processes, however, free nitrogen does take part; scientists have only recently found how to reproduce and use these processes. The utilization of the inert nitrogen of the atmosphere for agriculture, the basis of civilization and wealth, is one of the greatest triumphs of modern chemistry. The processes already developed and others which will be devised, no doubt, will be of incalculable value when the supply of nitrogen compounds available on the earth is exhausted or reduced to such an extent that it does not suffice to furnish the world's requirement. Inorganic compounds which contain nitrogen are soluble in water, and, as a consequence, we would not expect to find them accumulated in large quantities on the earth's surface except under unusual circumstances. In certain arid regions there are supplies of nitrates. The chief commercial source of these salts is the sodium nitrate obtained from Chile (Chile saltpeter) . Guano was formerly much used as an ingredient of fertilizers on account of the fact that it contains a large percentage of nitrogen compounds. Guano is obtained from certain tropical islands. It is the dried excrement of sea gulls, penguins, and other aquatic birds that breed in mil- lions along the coast of South America. Ammonia, NHs, and nitrates are found in fertile soils, and although they are present in but small proportions they are essential constituents; in this form they do not serve, however, as sources of nitrogen compounds for the industries. Coal contains compounds which yield ammonia on distillation; the manufacture of coke and illuminating gas by heating coal yields large amounts of ammonia. Gas-works, by-product coke ovens, and Chile saltpeter are the chief sources of the world's supply of combined nitrogen for industrial purposes. All living things contain nitrogen compounds; the so-called proteins, which are the chief constituents of flesh and are present in all vegetable matter, contain approximately 16 per cent of com- bined nitrogen. NITROGEN AND THE ATMOSPHERE 287 316. Preparation of Nitrogen. The gas in a comparatively pure condition can be readily obtained by burning a substance in air, and subsequently removing the products of combustion. This is done in the laboratory by burning phosphorus under a bell-jar, the bottom of which dips under water. Phosphorus is selected because it burns readily and the product of combustion, phos- phorus pentoxide, is a solid which dissolves in water. The process removes oxygen only and the gas left contains water-vapor, about 1 per cent of inert gases, and a trace of carbon dioxide and other substances. If desirable, all of these except the inert gases can be removed by the proper reagents from the nitrogen, but to exhibit the chemical inertness of the latter this is not necessary. A convenient way of removing oxygen from the air is to pass it over hot iron or copper, the latter being preferable in the laboratory because it is oxidized rapidly at a comparatively low temperature. On account of the expense of the materials used, nitrogen is not obtained in these ways for industrial purposes. When the oxygen is removed chemically, coke or coal is burned in the air, and the oxides of carbon produced are removed. Free nitrogen can be obtained from compounds of the element, but since the important commercial problem is to transfer nitrogen into these compounds, such processes are not used in chemical industry. All the oxides of nitrogen are reduced when heated with the copper or other metals; an example is illustrated by the following equation: 2NO + 2Cu = 2CuO + N 2 Ammonia can be oxidized by passing it over hot copper oxide: 2NH 3 + 3CuO = 3H 2 O + N 2 + 3Cu All organic compounds containing nitrogen yield nitrogen when heated with copper oxide. The decomposition is the basis for the quantitative analysis of such substances. The weight of the material burned and the volume of the nitrogen produced make it possible to calculate the percentage of nitrogen present in the com- pound. The easiest way to make pure nitrogen in the laboratory is to heat ammonium nitrite: NH 4 N0 2 = N 2 + 2H 2 288 INORGANIC CHEMISTRY FOR COLLEGES As the nitrite is a very unstable substance it cannot be kept. In making nitrogen in this way a strong solution of sodium nitrite and ammonium chloride is heated; the substances first interact by double decomposition and form ammonium nitrite, NH 4 C1 + NaNO 2 = NH 4 N0 2 + NaCl and the latter then decomposes into nitrogen and water. An important source of nitrogen for industrial purposes is liquid air. On evaporation of the liquid, nitrogen first escapes and then oxygen; liquid air serves, therefore, as a comparatively cheap source of both these gases. 317. Physical Properties of Nitrogen. Nitrogen, N 2 , is a color- less, odorless gas; it is slightly lighter than air (sp. gr. 0.967); 1 liter at and 760 mm. weighs 1.2506 grams; it is slightly soluble in water (2 volumes in 100 at 6). Liquid nitrogen boils at 194, and freezes at 214; since its critical temperature is 146 nitrogen cannot be liquefied by pressure at ordinary temper- atures; the liquid exerts such a high pressure that it is kept in open vessels. 318. Chemical Behavior of Nitrogen. At ordinary tempera- tures nitrogen is inert; the only chemical reaction known into which it enters under these circumstances is one brought about through the influence of bacteria present in nodules on the roots of certain plants, such as peas and beans and clover. The change of nitrogen to nitrates in this way is of great importance in agriculture and is the scientific basis for the rotation of crops. Under the influence of the bacteria some of the nitrogen of the air is con- verted into proteins and the nitrogen " fixed " in this way is changed by other bacteria present into nitric acid, which, penetrat- ing the soil, serves as a food for the growing plant. This subject has been much investigated, and the bacteria which affect the free nitrogen have been isolated. Seeds which have been dipped into a solution containing the bacteria inoculate the soil when they are sown, and the growth of the resulting plant is markedly increased. With rise in temperature the activity of nitrogen increases, and at very high temperatures it is one of the most active of all the elements. When an electric discharge passes through the air the nitrogen and oxygen in the immediate vicinity unite and form nitric NITROGEN AND THE ATMOSPHERE 289 oxide, NO. As a result of this action nitric acid, formed from this oxide, oxygen, and water-vapor, is produced in thunder storms. Nitrogen unites with hydrogen to form ammonia, NHs, when the two gases are heated together. Both of these reactions have been used for the so-called fixation of atmospheric nitrogen; nitric acid and ammonia, the two most important compounds of nitrogen, are now manufactured on the large scale from free nitrogen. At the temperature of the electric furnace nitrogen combines with certain metals and non-metals; some of the products, which are called nitrides, have interesting properties and may become articles of commerce. Some nitrides react with water to form ammonia. The reaction has been studied in the hope of finding a new method of making this industrially important substance. When magnesium is burned in the air the chief product is magnesium oxide, but at the temperature produced in the burn- ing, some of the metal unites with the nitrogen in the air and magnesium nitride, MgsN2, is formed. When the nitride is boiled with water, ammonia is produced. 319. Uses of Nitrogen. It has already been pointed out that free nitrogen is used in the preparation of nitric acid and ammonia, and that certain plants utilize the gas in their growth. An important use of nitrogen is in the manufacture of nitrogen- filled tungsten lamps. Up to a short time ago the bulbs of electric lamps contained no gas; the air was pumped out in order to prevent the action of oxygen on the filament. When tungsten replaced carbon as the material of which the filament was made, it was found that a black deposit formed on the walls of the bulb after the lamp had been used for some time. The deposit proved to be tungsten, which had distilled off from the filament at the high temperature to which it was heated. It was found that by filling the bulb with nitrogen the blackening was largely prevented. THE ATMOSPHERE 320. The components of the air that are present in almost constant proportions are oxygen, nitrogen, argon, and the other rare gases. Carbon dioxide is always present, but since it is the product of respiration, the combustion of coal and wood, and the 290 INORGANIC CHEMISTRY FOR COLLEGES decay of vegetable material, the quantity present in air varies greatly when the gas is introduced into the air from these sources. The amount of water-vapor in the air fluctuates between wide limits. Air is analyzed by determining the relative volumes of the gases present, and the results are expressed as percentages or as parts by volume. Air freed from water-vapor, carbon dioxide, and accidental constituents contains approximately 21 per cent oxygen, 78 per cent nitrogen, and slightly less than 1 per cent of argon. Samples of air collected at different parts of the earth's surface have been examined and the percentages of oxygen found never varied from one another more than 0.2 per cent. The actual amount of oxygen and other constituents in a sample of air varies, of course, greatly with the pressure of the atmosphere. The several constituents of the air will now be considered in some detail. 321. The Oxygen. The part that oxygen plays in natural processes has already been discussed (38). Life as we know it on the earth centers around oxygen. Whenever we move there is an expenditure of energy and this is supplied as the result of oxidation in the body. Our needs of this essential factor in life are well looked after. There is an ample, supply, and the body can func- tion when the proportion of oxygen in the air is much reduced. We can live in an atmosphere in which a candle will not burn. Oxygen converts animal and vegetable refuse through the agency of micro-organisms into innocuous gases. The amount of oxygen in air can be determined readily by introducing into a sample of known volume some substance which unites with oxygen at ordinary temperatures. The experiment is usually carried out by students in the laboratory with the aid of a eudiometer, which is a glass tube closed at one end and marked with lines etched into the glass so that the volume of the contained gas can be read. The open end of the tube is placed under water and a piece of phosphorus supported on the end of a wire is inserted into the air. The volume of the air in the tube is read, and the temperature of the outside air and the height of the barometer are recorded. After standing a day or longer, until all the oxygen has reacted with the phosphorus, the readings are made again. The decrease in volume equals the volume of the oxygen originally present. NITROGEN AND THE ATMOSPHERE 291 322. The Carbon Dioxide. By far the larger part of the carbon dioxide in the air is the product of the decay of vegetable material, and, as a consequence, the percentage of the gas in the atmosphere is practically constant, except in large cities or in the neighborhood of factories where great amounts of coal are burned. Country air contains 3 parts per 10,000, or .03 per cent; the atmosphere in a large city may have as much as 7 or 8 parts; and in a crowded room the content of carbon dioxide may rise as high as 50 parts. Air can be readily freed from carbon dioxide by passing it through a solution of sodium hydroxide. If it is desired to deter- mine whether carbon dioxide is present in a sample of gas, calcium hydroxide is used because in this case the carbonate formed is insoluble and precipitates; it is, therefore, visible. Barium hydroxide can be used in the test instead of calcium hydroxide. In making an air analysis, a large bottle (8 or 10 liters) is filled by blowing into it with a bellows, and thus obtaining a sample of the air in the neighborhood. A measured quantity of a solution of barium hydroxide of known strength is next added, and the bottle closed and shaken. After the precipitate has settled, some of the clear solution is drawn off in a pipette and the amount of barium hydroxide in it determined by neutralization with an acid. From the result obtained, the amount of the hydroxide which did not react with the carbon dioxide can be calculated; and the difference between this quantity and that used is a measure of the amount of carbon dioxide in the air analyzed. The decay of dead vegetation furnishes, as has been said, the bulk of the carbon dioxide in the air. This decomposition is brought about by oxidation induced by vegetable organisms which thrive in the presence of moisture. As a consequence, when we wish to preserve wooden buildings from decay we paint them. The oil contained in the paint protects the wood from the action of oxygen, and as moisture is not readily absorbed, the destructive organisms do not thrive. In time, however, the oil itself is oxidized, the paint crumbles, and the wood is exposed. Moisture plays an important part in this kind of decay; for this reason telegraph poles are often set in cement to keep the ends in the earth away from the water present in the soil. Wooden buildings last much longer in dry climates. Remains of temples built of wood hun- dreds of years old have been excavated in India. 292 INORGANIC CHEMISTRY FOR COLLEGES The burning of coal introduces large amounts of carbon dioxide into the air. It has been estimated that nearly one and one-half billion tons of coal are used each year. Forest fires also produce vast quantities of carbon dioxide, but these sources account for less than two-tenths per cent of the gas present in the atmosphere. Carbon dioxide gets into the air as the result of respiration in animals. Expired air contains about 4 per cent of carbon dioxide, and as a man breathes about 12 cubic meters per day, each indi- vidual on the earth is producing about 500 liters or 1000 grams of carbon dioxide daily. The constant accumulation of carbon dioxide in the air is prevented by the fact already emphasized, that growing plants convert the gas into the organic materials of which they are com- posed. Carbon dioxide is also removed from the atmosphere in rain and passes into rivers and finally into the ocean, which con- tains more of the gas than the atmosphere. Carbon dioxide in water is converted into carbonic acid and slowly decomposes mineral matter in the soil, and is in this way converted into soluble salts which are utilized by plants in their growth. 323. The Water-vapor. The fact that water is essential in the natural processes that take place on the earth has been empha- sized (96) . The distribution of water is brought about by the fact that it passes into the air as a gas and, consequently, becomes dis- seminated over the entire surface of the globe. It returns to the liquid state as rain or dew and thus furnishes plant life with one of its necessary foods. Water-vapor has been shown recently to play a very important part in ventilation. Various theories have been put forward as to the cause of the disagreeable sensations experienced in a crowded, ill-ventilated room; lack of oxygen, excess of carbon dioxide, and poisonous substances given off from the lungs have, in turn, been assigned as the cause. But none of these, as has been shown, is adequate. The changes in the amount of oxygen in the air do not appreciably affect us provided the required minimum is present; and it has been shown that a man can live comfortably when the amount of oxygen is reduced far below that obtainable in a room under ordinary circumstances. The amount of oxygen in the air on a mountain top is far below that in a crowded room but the lack causes no discomfort. The carbon dioxide in a room filled NITROGEN AND THE ATMOSPHERE 293 with people may increase to 50 parts in 10,000, but air containing this amount of the gas can be breathed without ill effects. The so-called " poisonous " material from the lungs is hypothetical, and its presence was assumed in searching for an explanation of the cause which produces the effects of bad ventilation. The body gives off substances of disagreeable odor and these produce, no doubt, marked psychological effects. On the other hand, experi- mentation has shown that the accumulation of water-vapor as the result of respiration from the lungs and skin in a crowded room has a marked effect on comfort. The air space between the indi- viduals, crowded together, soon becomes saturated with water- vapor, and the temperature rises. We have thus the conditions emphasized which prevail in the hot, muggy days of July; and the same discomfort is evident. Experiments have been made which prove the correctness of this theory; in one of them a num- ber of men remained in a small closet, tightly closed, until they suffered extreme discomfort from " lack of pure air." An elec- tric fan was then started; in a short time the moisture and other constituents of the air were evenly distributed throughout the closet and the experimenters lost the painful sensations they had experienced. Although carbon dioxide is not the cause of the effects pro- duced by bad ventilation, the amount present in the air is usually determined when a study of the ventilation in a room is being made. This is done because it has been shown by experience that the amount of carbon dioxide present is a rough measure of the state of the air; when it is high the room is uncomfortable, and if it is low the fact is evidence that a sufficient quantity of fresh air is being admitted. One of the simplest ways of removing the excess moisture and keeping a room at the proper temperature is to admit freely fresh air, and this method is the one used in proper ventilation. Electric fans are coming more into use and are val- uable aids. The facts stated above are the basis for a comparatively new science. The engineers who plan the ventilation of large audience chambers take into account, in estimating the amount of air to be admitted per minute, the number of people the room will accom- modate and the amount of oxygen consumed, carbon dioxide given off, and heat generated per individual. The humidity of the air 294 INORGANIC CHEMISTRY FOR COLLEGES supplied is also tinder control. In winter the amount of water- vapor in the air is small, because at low temperatures the vapor pressure of water is low. When such air is heated it is too dry for comfort; it must be brought into contact with water after being warmed, so that it can take up what is required to produce the relative humidity desired. It would lead too far to consider all these factors in detail, but the more important principles upon which ventilation is based are, no doubt, clear. 324. Other Constituents. Hydrogen peroxide is formed in minute quantities in the air. The way in which it is produced is not well understood, and there is opportunity for further study of the problem. It is possible that in natural processes it plays some part which has not been yet discovered. Hydrogen peroxide has been found in rain-water and snow; it may be produced as the result of electric disturbances in the atmosphere. Hydrogen peroxide is said to be formed by the action of sunlight on the surface of the ocean. It is produced in traces when zinc, lead, or copper rust in moist air. Hydrogen peroxide is unstable and soon decomposes in the presence of dust and bacteria. Ozone is present in the air after thunder-storms. The natural processes by which it is formed are, perhaps, similar to those which produce hydrogen peroxide. It is said to be a normal constituent of air near the sea and in forests, but adequate experi- mental evidence is lacking on this point. As far as is known it plays no significant part in nature. Hydrogen occurs in the air, but the amount present is exceed- ingly small, not more than 1 part in 1,500,000. The gas issues from volcanoes, and is formed as the result of putrefaction of organic matter brought about by certain bacteria. Nitric acid is formed from water and the oxides of nitrogen pro- duced as the result of electric discharges in thunder-storms. The acid dissolves in the rain and thus enters the soil and is a factor in supplying the combined nitrogen needed by plant life. The ammonia in the air is one of the products of the decay of organic material. It also is returned to the soil through the agency of the rain. It has been estimated that about 6 pounds of combined nitrogen per acre are obtained from the air per year. Sulphur dioxide gets into the air chiefly as the result of the burning of coal. In large cities its presence is of importance, for NITROGEN AND THE ATMOSPHERE 295 it is converted into sulphuric acid, which attacks materials of construction made of metal if they are not well protected. The effect is generally noticeable near large railroad stations. 325. Dust is found in all air under normal conditions. Its presence in the atmosphere is the cause of the formation of clouds and fog. Water- vapor will not change into a liquid unless it comes in contact with a surface upon which it can condense. When moist air is cooled to the temperature at which it is sat- urated with water-vapor, condensation takes place on the par- ticles of dust and a cloud is formed. If the temperature continues to fall, the minute drops of water in the cloud grow larger as the result of the condensation of more water, and finally reach such a size that they can no longer stay suspended and fall as drops of rain. The organic dust in the air contains micro-organisms such as bacteria and spores of fungi and molds. These several organisms bring about different kinds of fermentation; one causes the forma- tion of wine from grape juice as the result of the production of alcohol, and another converts cider into vinegar, which contains acetic acid. Putrefactive bacteria decompose animal material, and poisonous products and gases with a foul odor are produced. The air also carries at times pathogenic bacteria which serve to transmit disease. Many of the organisms in the air thrive best in damp places away from direct sunlight. This is one reason for having health resorts at high altitudes in a sunny climate. Air can be readily freed from dust, both mineral and organic, by being drawn through a layer of raw cotton. The medium in which bacteria are grown in experimental work is protected from organisms of the air by placing a bit of cotton wool in the mouth of the test-tube containing the material under investigation. The presence of argon and helium in the air has been men- tioned. These gases are of such interest and their study has led to such important results, that they deserve a somewhat detailed consideration. ARGON AND RELATED GASES 326. In describing argon and the other gases like it in the air it seems advisable to treat the subject historically and outline the researches which led to such interesting results. In this way the 296 INORGANIC CHEMISTRY FOR COLLEGES student will become acquainted with the spirit which guides scien- tific research and the methods used in the attack of a problem. Lord Rayleigh, a physicist, undertook to determine the density of certain gases with greater accuracy than had been attained before. When a property of -a substance is to be determined accurately it is necessary, of course, to have it in the purest con- dition possible. In order to avoid the possibility of the presence of unknown accidental impurities which may not be removed by the processes of purification used, it is advisable to study samples of the substance obtained from widely different sources in which the accidental impurities may be different. Lord Rayleigh took the precaution to weigh with great accu- racy nitrogen obtained from the air and that formed as the result of the decomposition of certain compounds containing the element. The atmospheric nitrogen when freed from all known substances was weighed in a bulb which contained 0.14332 gram of the gas measured under standard conditions. When the same bulb was filled with nitrogen obtained from ammonium nitrite, the gas under the same conditions weighed 0.14256 gram. This difference was very small, but considerably more than the experimental error, and it was always found as the result of repeated experiments. In the case of the other gases studied similar results were not obtained; the weight of each gas was constant whatever its source. 327. The next step was to explain the facts, and as the ques- tion was a chemical one Professor Ramsay, a chemist, was asked to co-operate in the investigation. Various hypotheses were put forward: One or the other or both samples of nitrogen might be impure; the atmospheric nitrogen might not have been freed completely from known substances or it might contain a new sub- stance, heavier than pure nitrogen, which was not removed in the processes used; or the nitrogen from the ammonium nitrite used in the preparation of the gas might contain a trace of hydrogen. These suggestions were tested one after another, and it was found that all known substances had been eliminated; the logical con- clusion was that the air contained a very small amount of an unknown gas heavier than nitrogen. The next problem was to isolate the gas. To do this nitrogen containing the unknown substance was repeatedly passed over NITROGEN AND THE ATMOSPHERE 297 heated magnesium which united with the nitrogen and therefore removed it. The residual gas had a volume about 1 per cent of that of the nitrogen used. Nitrogen was also removed by a method that had been used by Cavendish in 1785. The gas was sparked in the presence of oxygen; under these circumstances the nitrogen was converted into nitric acid, which was removed by dissolving it in a solution of sodium hydroxide. The excess of oxygen was taken up by passing the gas over hot copper. It is worthy of note here that when Cavendish made his experiment he found that he always obtained a residue which could not be removed by continued sparking, and he was unable to explain his results. He states that the volume of the residual gas was T|-O of the nitrogen used a figure which agrees well with the results of the experiments carried out over a century later. The properties of the gas obtained in this way were care- fully studied. It was found to be twenty times as heavy as hydrogen, whereas nitrogen is fourteen times as heavy. This fact accounted for the difference in weight of the nitrogen sam- ples observed by Rayleigh. 328. The gas was evidently a very stable substance, since it had resisted very active chemical reagents and high temperatures in the course of its isolation. It was thought it might have been formed from the constituents of the air as the result of these vig- orous processes; it might, for example, be a polymer of nitrogen or some unknown compound. To test this view nitrogen was freed from known gases at a low temperature and allowed to pass slowly through unglazed clay pipes surrounded by a vacuum. It will be recalled that heavy gases diffuse more slowly than light ones (50) . The gas was passed repeatedly through the clay pipes and each time a part of it was lost by diffusion through the walls of the pipes into the vacuum. Samples of the gas were weighed from time to time and it was found that they increased in weight; the lighter nitrogen diffused more rapidly than the heavier gas, which, accordingly, became more concentrated. It was evident from these experiments that the new gas was present in the air and was not formed as the result of any chemical reaction taking place in the process of its isolation. 329. The chemical properties of the gas were then studied and it was found that it entered into chemical reaction with nothing. 298 INORGANIC CHEMISTRY FOR COLLEGES It was subjected to the most active reagents; metals, sodium hydroxide, sodium peroxide, sodium and calcium persulphides, all red hot, did not affect the gas. Nascent chlorine and fluorine, the most active of all the elements, failed to affect it. It with- stood the silent electric discharge and the electric arc. There was no known element that would not enter into chemical reaction under many of the conditions to which the new gas was subjected. This unique behavior led to the selection of the name argon for the gas, the word being derived from the Greek word signifying inactive or idle. The atomic weight of argon was found to be 40 and it was shown to have one atom only in the molecule. 330. No use was found for argon for a number of years, but its inertness and the fact that it is a monatomic gas finally led to an important application. The part that nitrogen plays in the gas-filled tungsten lamp has already been explained (319). While the presence of nitrogen prevents the blackening of the globe by the material volatilized from the filament, it was found that the lamp grew hot and, as a result, an undue proportion of the elec- trical energy was converted into heat. The amount of heat lost from the lamp depended, evidently, on the rate at which the heat was conducted away from the filament. Heavy molecules in the gaseous condition move more slowly than light ones (50); as a consequence, the rates at which gases conduct heat vary with the densities of the gases. The replacement of nitrogen, which has the molecular weight 28, by argon (40) reduced the rate at which heat energy was lost from the lamp. Another reason for the re- placement of nitrogen by argon is based on the fact that the heat required to raise to the same extent the temperatures of equal volumes of the two gases, is much less in the case of argon. The difference is traceable to the monatomic structure of the inert gas. Argon was substituted for nitrogen and a more efficient lamp was the result that is, a greater proportion of electrical energy was transformed into light because less heat was lost. The argon required was obtained from liquid air. We shall have more to say of these lamps later when tungsten is considered. They represent the combination of the results of many researches in widely different fields researches inspired by different motives. Rayleigh's desire to weigh nitrogen accurately, NITROGEN AND THE ATMOSPHERE 299 Scheele's study of the mineral now known as scheelite, which resulted hi the discovery of tungsten, and the work of the physi- cists who devised the way to liquefy air, produced results which were available to the inventors and engineers who produced the most efficient means of artificial illumination. The researches mentioned were in what is sometimes inappropriately called pure science; they were carried out with the aim of finding out the facts and without thought of their application. The scientists who developed the lamp had a definite object in view; their researches were in what is called applied science. They utilized the results obtained by many other workers, and selecting here and there, attained the desired result. Incidentally important new facts and principles were discovered that may prove valuable later. The methods of research are the same in all fields; in applied science, however, the aim is very definite some particular thing is to be done and the problem is oftentimes for this reason a more difficult one to solve than a problem in pure science. It is impor- tant to note, however, that the principles and facts of pure science are the building blocks of applied science. In order to do a thing better than it has been done in the past, or to do a new thing, we must be able to call to our aid knowledge that has not been utilized for this purpose in the past. And it often happens that things considered highly theoretical or impractical have become of vital importance in the production of things of practical value. The more one knows of facts and principles which have not been applied, the greater the store to be drawn on when new problems arise. 331. Helium. In seeking a source of argon other than the air, Ramsay was led to examine certain minerals which were reported to yield free nitrogen when heated. Among these was the mineral clevite. The gas given off was examined with a spec- troscope and found to be a mixture in which the element helium was present. Lockyer in 1863, in examining the spectrum of the chromosphere of the sun during an eclipse, noted an orange line which had not been observed in the spectra of anything of terres- trial origin. He drew the conclusion that the sun contains *an element unknown on the earth, and he called it helium, He, deriving the name from the Greek word meaning the sun. Ramsay studied the gas in the way used with argon and found 300 INORGANIC CHEMISTRY FOR COLLEGES that it resembled it in being inactive and monatomic. It was, however, a very light gas, being only twice as heavy as hydrogen and one-tenth as heavy as argon; its atomic weight was found to be four. The world was searched over by many investigators for other sources of helium and it was discovered occluded in a number of minerals, dissolved in the water of certain mineral springs, and as a constituent of natural gas from certain sources. It was found later to be present in the air to the extent of 1.4 parts per million. Helium was liquefied by Onnes by subjecting it to the temperature obtained by boiling liquid hydrogen in a vacuum. It has a lower boiling-point than that of any other known substance (268.5 or 4.5 Abs.). When helium was boiled in a vacuum a part solidified at a temperature which was estimated to be 2 absolute. The most interesting fact about helium is that it is found in the gaseous emanation given off by radium, and is a product of the spontaneous decomposition of the element. This far-reaching fact will be discussed in some detail later, since it has led to a new conception of the constitution of matter and markedly enlarged our knowledge of the genesis of the earth and the materials of which it is composed, 332. We have in helium another striking example of how a sub- stance of theoretical significance suddenly becomes of great prac- tical interest. Again, it is the combination of certain properties that make it valuable. In the case of helium it is its extreme lightness and chemical inactivity. It was suggested during the recent war that helium could be substituted for hydrogen in dirigible and observation balloons. The gas has almost the same lifting power as hydrogen, and being non-inflammable could not be ignited by tracer bullets which were invented during the war to carry a flame to a balloon and explode the hydrogen which it contained. Natural gas obtained in certain parts of Texas con- tains 0.85 per cent of helium. The advisability of using helium for war purposes was so great that the problem of extracting it from a source in which it was present in such small quantity was undertaken. The process adopted involved liquefaction of the nakiral gas and isolation of the helium from the residue which did not liquefy. The pure gas compressed in steel cylinders was ready for transportation when hostilities ceased. There is no doubt that this use for helium will be developed, since great dan- NITROGEN AND THE 'ATMOSPHERE 301 ger is avoided by replacing hydrogen in balloons by a non-inflam- mable gas. 333. Neon, Krypton, and Xenon. Ramsay continued his search for other gases and distilled a large amount of liquid argon obtained from the air. He found that he had a mixture, and in addition to pure argon small quantities of helium and of three new gases were obtained. These resembled argon in being chem- ically inert and monatomic. Greek words were again used in naming the gases. Neon (new) has the atomic weight 20.2, krypton (hidden) 82.9, and xenon (stranger) 130.2. Krypton is of interest because the lines of its spectrum have been observed in the aurora borealis or northern lights. No adequate explana- tion of this striking natural phenomenon has yet been given and there is no reason known why krypton should be present in quan- tities sufficient to make its spectrum visible. The number of volumes of air which contain one volume of the inert gases are as follows: Argon, 106.8; neon, 80,800; helium, 245,000; krypton, 20,000,000; xenon, 170,000,000. EXERCISES 1. (a) Calculate the weight of 1 liter of He. (6) Calculate the relative lifting power of helium and hydrogen when used in balloons. 2. (a) How could you determine the amount of air exhaled per day by a man at rest? (6) Make the experiment in the laboratory and calculate the amount of carbon dioxide produced in twenty-four hours assuming that exhaled air contains 4 per cent CO 2 . 3. A sample of air was analyzed for carbon dioxide with the following results: 10 liters of air were shaken in a bottle with 100 c.c. of 0.1N Ba(OH) 2 . After the precipitated BaCO 3 had settled 50 c.c. of the clear liquid was drawn off and neutralized with a 0.1N solution of an acid, 29.10 c.c. of the solution being required. Calculate the (a) weight of CO 2 in the sample of air, (6) percentage by volume of CO 2 in the air, and (c) the parts CO 2 in 1000 by volume. 4. The solubility of O 2 in water at and 76 cm. is approximately 4 vol- umes in 100 volumes, (a) Assuming air to be made up of f N 2 and 1 O 2 , what is the pressure of the N 2 and O 2 in air? (&) What volume of the two gases would be dissolved in 100 c.c. of water if air were bubbled through the latter until it was saturated? (See Henry's law.) (c) How could you determine by an experiment the relation between the N 2 and O 2 dissolved by water when in contact with air? (d) If air were a compound of nitrogen and oxygen what relation would exist between the nitrogen and oxygen in the air and in the part soluble in water? 5. State two reasons for the belief that air is a mixture and not a com- pound. CHAPTER XXII AMMONIA AND ITS DERIVATIVES 334. The chemistry of ammonia is of the greatest importance because of its significance in the growth and decay of living things. The farmer fertilizes his fields with ammonium salts and later when he eats the grain or vegetables he has grown he obtains the nitrogenous material which through digestion and assimilation becomes a part of his flesh. Since the supply of ammonia fur- nished by nature is limited, the chemist has devised ways of con- verting the nitrogen of the air into this important product, and has triumphed over what appeared at first insurmountable difficulties in his endeavor to make out of the inexhaustible supply of air and water a substance that could be used as a plant food when the naturally occurring nitrogen fertilizers are exhausted. Since ammonia is produced as the result of the decay of refuse organic matter, its presence in undue amounts in natural water is evidence of contamination. As a consequence, in examining a water supply to be used for drinking purposes, the amount of ammonia present is always determined; the results obtained guide the chemist in his decision as to its potability and may lead to the discovery of sources of contamination which can be removed. The physics of ammonia serves mankind in another important way. The gas can be readily liquefied through pressure alone and the resulting liquid, boiling at a very low temperature, absorbs a large amount of heat when it passes into gas its latent heat of vaporization is very high. These facts have been utilized in making machines for the production of low temperatures (180). By the use of ammonia we can make ice anywhere in unlimited quantity, or keep a warehouse cold for the storage of perishable food materials. This application of ammonia has had a marked effect on the health of people and has revolutionized the economics of food production; refrigerator cars and cold-storage warehouses 302 AMMONIA AND ITS DERIVATIVES 303 have become a necessity in modern civilization with the growth of large cities. 335. History and Occurrence of Ammonia. Ammonia is formed when animal matter decays, and as it has a characteristic odor it has been known from the earliest times. Ammonium chloride the compound formed from ammonia and hydrochloric acid was known to the Egyptians. Its ancient name, sal ammo- niac, from which the word ammonia is derived, is said to have been given to the salt in honor of the Egyptian sun god Ammon. In the Middle Ages ammonia gas was made by distilling the horns of harts, and was supposed to be very valuable as a medicinal agent. It was called spirits of hartshorn, a name by which it was known until very recently when used in medicine. The alchemists obtained ammonium chloride by evaporating to dryness and heating to a high temperature a mixture of urine and salt. Priestley, in 1774, first isolated ammonia as a gas as the result of heating ammonium chloride with lime. He collected the gas over mercury. Many of Priestley's most important discoveries were due to the fact that he used mercury in his pneumatic trough instead of water. He isolated in this way for the first time, in addition to ammonia, sulphur dioxide and hydrogen chloride, gases which could not have been recognized if water had been used. Priestley's most important discovery that of oxygen was also made in his pneumatic trough and gas tube filled with mercury. It has often happened that important discoveries have sprung from the use of new methods of carrying out chemical operations or reactions. One can never foretell what may result from such a new method; and the discoverer has added a new tool to the science. 336. Preparation of Ammonia. The commercial source of ammonia is ammonium salts, and the gas is prepared from these by the action of a base; ammonium chloride and calcium hydrox- ide are commonly used in the laboratory. As the result of the reaction, which is one of double decomposition, ammonium hydroxide is first formed and then decomposes into ammonia and water. The equations for the reactions are as follows: 2NH 4 C1 + Ca(OH) 2 = 2NH 4 OH + CaCl 2 NH 4 OH = NH 3 + H 2 304 INORGANIC CHEMISTRY FOR COLLEGES The reaction is brought about with the solid substances containing only a trace of water, on account of the fact that ammonia is very soluble in water. The gas generated is collected in the laboratory by downward displacement of air. The collecting vessel is held with the mouth downward because ammonia is only slightly more than half as heavy as air and, consequently, rises. When calcium hydroxide reacts with ammonium chloride in aqueous solution the double decomposition represented by the equation given above is brought about as the result of the fact that ammonium hydroxide breaks down spontaneously to form ammonia. In general, ammonia is formed when a substance that furnishes an ammonium ion is treated in solution with a substance that yields a hydroxyl ion. The chief sources of ammonia are the products obtained as the result of the distillation of coal to make illuminating gas or coke for metallurgical uses (226). The gases formed in the dis- tillation are scrubbed with water. The solution which results is heated with lime to decompose the ammonium salts formed as the result of the neutralization of a part of the ammonia by the carbon dioxide and other acidic material produced. The gas which escapes is absorbed in sulphuric acid and the resulting solution on evapora- tion yields ammonium sulphate. Hydrochloric acid is at times used instead of sulphuric acid and ammonium chloride is obtained. Coal contains only about 1 per cent of nitrogen, and the yield of ammonia is, consequently, small, but such large quan- tities of coal are distilled that the supply of ammonia has been adequate. When coke is made in bee-hive ovens (224), all the volatile materials are lost. As the demand for ammonia increased, more and more of these ovens were replaced by by-product ovens in which these valuable materials are condensed and saved. 337. Synthetic Ammonia. It has been known for many years that ammonia is formed when electric sparks are passed through a mixture of nitrogen and hydrogen. The behavior of nitrogen with hydrogen under these circumstances is quite different from that of oxygen. A single spark is all that is required to cause the union of the entire amount of oxygen provided enough hydrogen is present; the reaction propagates itself through the entire mass. We could synthesize water from oxygen and hydrogen in any desired quantity if it were necessary. With nitrogen, however, AMMONIA AND ITS DERIVATIVES 305 only the gases in the immediate vicinity of the flame unite and the mixture has to be sparked for a long time to get appreciable quantities of ammonia. 338. The reaction between nitrogen and hydrogen has been studied exhaustively with the hope of working out a method for the synthesis of ammonia which would be of commercial value. Ammonia itself is a necessity, but the fact that it can be readily oxidized to nitric acid makes its commercial synthesis of added importance. The supply of nitric acid is largely obtained from Chile saltpeter, and as this is not inexhaustible chemists have been studying in recent years the " fixation of atmospheric nitrogen." Since nitric acid is involved in making the explosives used in warfare, and as the sodium salt is an important ingredient of fer- tilizers, the governments of several nations have recently sub- sidized the study of this important problem. During the recent war the blockade of Germany prevented the importation of salt- peter from Chile and, as a consequence, large amounts of ammonia were synthesized and converted into nitric acid. The Germans, in all probability, had foreseen the possibility of a blockade in case of war and had bent their energies to devise a commercial synthetic method for the preparation of ammonia. Nitric acid, we shall soon see, can be made directly from the air and water, but the method used requires large amounts of electrical energy and can be employed only where the energy is cheap as the result of the availability of water-power. The production of nitric acid from ammonia and the synthesis of the latter itself do not require electrical energy, and, therefore, this method is preferred where elec- tricity must be made by burning coal. The pressing need of Germany was finally solved by Haber, and the process of making ammonia from nitrogen and hydrogen is known by his name. 339. We shall first examine carefully the reaction by which ammonia is formed from nitrogen and hydrogen, and then see how it was put on a manufacturing basis. The study of this case will bring out clearly how the chemist is able to modify the con- ditions under which a reaction takes place and thus make it pos- sible to obtain the desired result. In the application of a chemical reaction to the preparation of a substance it is necessary to know first whether the reaction is a reversible one. If this is the case the equilibrium attained in the 306 INORGANIC CHEMISTRY FOR COLLEGES reaction under different conditions must be determined in order to get the largest possible amount of the desired product. The equilibrium in a reversible reaction is affected by the temperature and, in many cases, by pressure, and the rate at which it is reached is modified by catalyzers. If a chemical reaction involving gases is to be intelligently utilized it must be studied from the above point of view and a large amount of data accumulated, by a con- sideration of which the most efficient conditions can be arrived at. The reaction by which ammonia is synthesized is as follows: N 2 + 3H 2 * 2NH 3 + 2 X 12,200 cals. The amount of heat developed must be determined because the reaction is a reversible one and the equilibrium attained is deter- mined by the temperature. In planning the apparatus in which the reaction is to be carried out, provision must be made to keep the materials at the proper temperature; this is done by with- drawing a part of the heat formed or by adding more heat as the case requires. The equilibria with change in temperature must be deter- mined, and since the reaction is exothermic in this case, rise in temperature is accompanied by shifting of the equilibrium to pro- duce a greater percentage of nitrogen and hydrogen (van't HofFs law, 267). The reaction must then be run at as low a temperature as is consistent with the production of a practical amount of ammonia. The lowering of the temperature has the undesired effect of slowing up the rate of reaction. This result was overcome by finding a suitable catalyzer to increase the rate. It is seen from the equation that the gases react in the relation of 1 molecule of nitrogen to 3 of hydrogen and form 2 molecules of ammonia. We have seen (77) that this indicates that 1 volume of nitrogen unites with 3 volumes of hydrogen and forms 2 volumes of ammonia; there is a contraction from 4 volumes to 2. When there is a change in volume as the result of a chemical reaction the pressure under which it takes place has an effect on the equilibrium (law of Le Chatelier, 269). The result is what we would expect increased pressure favors a reaction which leads to decreased volume. It is evident that a larger percentage of nitrogen and hydrogen is converted into ammonia as the pressure under which the reaction takes place is increased; the amount of pressure to AMMONIA AND ITS DERIVATIVES 307 be used is determined by the engineering difficulties encountered in the construction of the apparatus. In order to determine the best conditions under which to bring about the reaction, a large amount of data had to be accumulated from accurate measurements of the factors involved, and an exhaus- tive search made for a proper catalyzer. It was found that the best practical conditions were a temperature of 450, a pressure of 200 atmospheres, 1 and a catalyzer of iron which contained small quantities of other substances the nature of which is a trade secret. To fulfill these conditions was a difficult engineering feat; at the high temperature required hydrogen passes slowly through most metals, and their strength decreases. To construct an apparatus of the size required to resist such a high pressure was, therefore, difficult. By using steel of a special composition and by lining the apparatus with an alloy that resisted the action of hydrogen at the temperature used, the problem was finally solved. During the course of the work, however, explosions of great violence occurred and great steel tubes, which contained the catalyzer, and were 50 feet long, 3 feet in diameter, and 9 inches thick were blown to pieces. 340. The production of the nitrogen and the hydrogen for the synthesis presented a problem which was solved in an equally thorough and scientific way. When air is passed over hot coke the nitrogen is unaffected and the oxygen converted into either carbon dioxide or carbon monoxide or a mixture of the two, the temperature and other conditions determining the composition of the resulting mixture (200, 211) : C + O 2 = CO 2 CO 2 + C + 2CO If the oxides are removed from the product, the reaction furnishes a source of nitrogen. When steam is passed over hot coke, hydrogen and carbon monoxide and carbon dioxide are formed (228) : H 2 + C = H 2 + CO CO + H 2 O < CO 2 + H 2 1 It has been reported recently that Claude, a Frenchman, has constructed apparatus in which the reaction can be carried out at 1000 atmospheres. This change would increase the percentage of ammonia in the gases at equilibrium. 308 INORGANIC CHEMISTRY FOR COLLEGES If the oxides of carbon are removed in this case the reaction fur- nishes a source of hydrogen. By passing air and steam together over hot coke the resulting gas is a mixture of nitrogen, hydrogen, and the two oxides of carbon. The proportions of the several constituents are determined by the relative amounts of air and steam used and the temperature at which the reaction is carried out. Since carbon monoxide is difficult to remove from a gas on account of its lack of solubility in common liquids, and carbon dioxide can be readily removed, it is necessary to run the reaction under the conditions which yield the highest attainable percentage of the dioxide. These condi- tions can be determined by a study of the equilibria in the revers- ible reactions noted above. As it was not possible to convert all the carbon monoxide into dioxide in a single operation, the mixture of gases first obtained was mixed with a fresh supply of steam and passed over a catalyzer which increased the rate of the reaction CO + H^O <= C02 + Eb. The catalyzer in this case is a special variety of iron oxide. The reaction just given is a reversible one, so the carbon monoxide was not completely removed. The trace that remained interfered with the union of nitrogen and hydrogen and was dissolved out by passing the gas through a solution containing a cuprous salt (214). To remove the large amount of carbon dioxide in the gases by means of lime was impracticable on account of the expense, and water was used. The solubility of a gas in water increases rapidly with increased pressure (Henry's law, 185), and as the gas had to be compressed for the final synthesis it did not add to the cost to take advantage of the increased solubility of carbon dioxide under these condi- tions. By washing it at 50 atmospheres pressure with water, the carbon dioxide was, accordingly, removed. A large part of the energy of the compressed carbon dioxide was utilized by allowing it to escape from the water solution into a turbine engine which furnished power. 341. To sum up, the Haber process for making synthetic ammonia is as follows: Nitrogen and hydrogen are obtained by passing air and steam over hot coal or coke; the resulting gases, mixed with more steam, are passed over a catalyzer in the presence of which the carbon monoxide present, by reacting with the water, is converted into carbon dioxide. The gases are compressed to AMMONIA AND ITS DERIVATIVES 309 50 atmospheres and brought into contact with water which dissolves out the carbon dioxide. They are next passed at this pressure through a solution of cuprous formate in ammonia which removes the trace of carbon monoxide remaining. After com- pression to 200 atmospheres the mixture of nitrogen and hydro- gen is passed over a catalyzer and unites to form ammonia. Since under the most favorable conditions only about 2 per cent of the gases react, they are passed, still under the high pressure, through water which absorbs the ammonia, and are then returned to the catalyzer chambers. The circulation of the gases is con- tinued and additional nitrogen and hydrogen admitted at the rate at which they unite. The water containing the ammonia is drawn off when it contains enough of the gas to saturate it at atmospheric pressure and temperature. The ammonia is used as such or converted into nitric acid in a way to be described later. The Haber plant cost about $25,000,000 and about seventy-five chemists and physicists are engaged in the control of the process and in doing the research work required for its development. The production of nitric acid manufactured from the ammonia made in this way was approximately 100,000 tons per year. The synthesis of ammonia by the Haber process has been described in some detail because it is an excellent illustration of how a solution of an important industrial chemical problem has been reached through the application of physical chemistry. It emphasizes the importance of the study of chemical equilibrium and the factors which govern it, all of which are susceptible of mathematical analysis. We have seen how van't HofFs law, Henry's law, and the principle of Le Chatelier, have been applied to the real service of mankind. The development of industrial chemistry in all its phases is rapidly taking place as the result of the application of the fundamental principles underlying the science, which are considered fully in physical or theoretical chemistry. 342. The Cyanamide Process for Ammonia. Where elec- trical power is cheap this process for the production of ammonia can be advantageously used. Calcium cyanamide decomposes slowly with water and yields ammonia, and for this reason is used itself as a fertilizer, as it furnishes a cheap source of this important plant food. The compound has been made in large 310 INORGANIC CHEMISTRY FOR COLLEGES quantities in Norway and in the United States where water power for generating electricity is available. In Germany brown coal and lignite are abundant, and furnish a cheap source of power. When nitrogen is passed over calcium carbide (217) at the tem- perature of an electric furnace, the gas is absorbed and calcium cyanamide is formed : CaC 2 + N 2 = CaCN 2 + C The compound decomposes slowly with water with the formation of ammonia: CaCN 2 + 3H 2 O = CaCO 3 + 2NH 3 If steam is used the decomposition is rapid and by condensing the vapor a solution of ammonia is obtained. In calcium cyanamide the calcium is in combination with the nitrogen. It will be recalled (318) that compounds of metals with nitrogen are decomposed by water, and that the hydrogen of the latter unites with the nitrogen to form ammonia, and the oxygen with the metal to form an oxide. The reaction of calcium cyana- mide with water takes place in an analogous way; in this case, however, the carbon in the compound is converted into carbon dioxide which reacts with the calcium oxide and forms calcium carbonate. 343. The formation of ammonia when animal matter undergoes decomposition has been emphasized. An example of this is the decomposition of urine, which results in the conversion of the urea present into ammonia and carbon dioxide, through the agency of a micro-organism present in the air. Ammonia is formed when the proteins and other nitrogenous materials in plants and animals are heated with concentrated sul- phuric acid. The acid oxidizes these compounds, and the carbon they contain is changed into carbon dioxide, the hydrogen into water, and the nitrogen into ammonia. The latter unites with the excess of sulphuric acid to form ammonium sulphate. Advantage is taken of this fact in analyzing quantitatively food products for nitrogen. Proteins are essentials in food and their determination in the valuation of a food product is, therefore, important. The product to be analyzed is heated with concentrated sulphuric acid, and when the oxidation is complete the solution is made AMMONIA AND ITS DERIVATIVES 311 alkaline with sodium hydroxide and heated. The ammonia set free distills over with water; its amount is determined by neu- tralizing it with an acid of known concentration. The procedure outlined is used in what is called the Kjeldahl method for the deter- mination of nitrogen. The oxidation with sulphuric acid is cata- lyzed by mercury salts. 344. Physical Properties of Ammonia. Ammonia is a colorless gas which is lighter than air (sp. gr. 0.5971). Its critical tempera- ture is 131 and it can, therefore, be condensed to a liquid at ordinary temperatures. Liquid ammonia boils at 33.5 and solidifies to a white crystalline solid at -77. Its molecular heat of vaporization is 5700 calories, the value being greater than that of any other liquid except water. Ammonia is very soluble in water; 1 volume of water at dissolves 1300 volumes of the gas, and at 16, 783 volumes. The concentrated ammonia of commerce contains about 28 per cent by weight of the gas and has the specific gravity 0.9. Liquid ammonia is an excellent solvent for both inorganic and organic substances. It dissolves sodium and potassium and other substances which are insoluble in or react with water. When salts dissolve in liquid ammonia they conduct the electric current and show the characteristic behavior exhibited when ionized in water. The liquid is transported in iron cylinders and is used in ice- machines (180). The pressure in a cylinder of liquid ammonia at 20 (68 F.) is about 8 atmospheres, and at 30 (86 F.) about 11.5 atmospheres. 345. Chemical Behavior of Ammonia. The heat of formation of ammonia is small compared with that of water and we would expect, therefore, to find it less stable toward heat. When 1 gram- molecule (17 grams) of ammonia is formed from nitrogen and hydrogen 12,200 calories are set free; the production of a corre- sponding amount of water- vapor (18 grams) yields 58,100 calories. The gas is almost completely decomposed into nitrogen and hydrogen at 700. The change in volume which occurs when ammonia decomposes can be determined by passing electric sparks through a measured volume of the gas collected in a glass tube over mercury. When the sparking no longer produces a change it will be found that the volume of the gases produced is twice that of the ammonia taken. 312 INORGANIC CHEMISTRY FOR COLLEGES The relative proportions of the hydrogen and nitrogen in the resulting gases can be determined by passing them over hot copper oxide, which converts the hydrogen into water. The results show that the volume of hydrogen is three times that of the nitrogen. The equation for the reaction is 2NH 3 <= N 2 + 3H 2 When ammonia is heated with the more active metals, such as magnesium, nitrides are formed : 3Mg + 2NH 3 = Mg 3 N 2 + 3H 2 At lower temperatures with sodium, only a part of the hydrogen is set free and sodium amide results : 2Na + 2NH 3 - 2NaNH 2 + H 2 Ammonia forms addition-products with certain salts, which resemble in composition the hydrates of these salts; for example, compounds of the composition CaCl 2 ,6NH 3 and CaCl 2 ,6H 2 O, and CuSO4,5NH 3 and CuSO4,5H 2 are known. It is evident that anhydrous salts which unite with ammonia cannot be used for drying ammonia gas containing water-vapor. The solubility in water of the compounds of salts with ammonia, is often quite different from the solubility of these salts themselves. For example, silver chloride, AgCl, is almost insoluble in water, whereas its compound with ammonia, AgCl,2NH 3 , is readily soluble in water. It is for this reason that precipitated silver chloride dis- solves in a solution of ammonia. The change in solubility when many inorganic substances are treated with ammonia is made use of in qualitative analysis. Ammonia does not burn when a stream of the gas issuing from a tube is brought into contact with a flame. If surrounded by an atmosphere of oxygen, however, it burns with a yellow flame on ignition. This can be easily demonstrated by passing oxygen to the bottom of a vessel containing a concentrated aqueous solution of the gas; as the oxygen rises through the liquid it carries along with it some ammonia, and the gases when they leave the solution can be ignited; they burn with a yellow flame. The product AMMONIA AND ITS DERIVATIVES 313 formed depends upon the proportion of oxygen; when an excess of ammonia is present nitrogen is formed : 4NH 3 + 3O 2 = 2N 2 -f 6H 2 O With more oxygen, oxides of nitrogen are produced. By burning ammonia with air in the presence of a catalyzer nitric oxide can be formed : 4NH 3 + 5O 2 = 4NO + 6H 2 O We shall see later that nitric oxide unites with oxygen from the air in the presence of water to form nitric acid. The reaction is the basis of an important new chemical industry which will be described in the next chapter. An aqueous solution of ammonia shows all the characteristic properties of a solution of a base. For this reason it was believed for many years that ammonia reacted with water to form a com- pound of the composition NH-iOH, although no one had suc- ceeded in isolating it. Very recently this compound, ammonium hydroxide, has been isolated at low temperatures; it forms crystals which melt at 79. 346. The reaction between ammonia and water is a reversible one: NH 3 + H 2 O ^> NH 4 OH It has been shown that in a molar solution (17 grams of NH 3 in 1 liter of solution) at 20 about 30 per cent of the ammonia is in combination with water as ammonium hydroxide. As the tem- perature rises the decomposition indicated by the equation when read from right to left takes place rapidly, and at the boiling- point of water it is practically complete. Water can be com- pletely freed from ammonia by boiling it. Ammonium hydroxide is a comparatively weak base, that is, the extent to which it ionizes in water is much less than is the case with sodium hydroxide; it ionizes as indicated below: NH 4 OH = NH 4 + In a molar solution of sodium hydroxide 72 per cent of the base is in the form of ions, whereas in a molar solution of ammonia only 0.4 per cent is ionized. In solutions of one-tenth this concen- 314 INORGANIC CHEMISTRY FOR COLLEGES tration the ionization values are 91 per cent and 1.3 per cent, respect- ively. The caustic properties of a base are due to the hydroxyl ions which it produces. Sodium hydroxide, for example, decomposes and dissolves protein material, and for this reason when a solution of it is allowed to stay in contact with the flesh a painful " burn " results; with ammonium hydroxide, however, this does not occur because the concentration of the hydroxyl ions in the solution is so small. 347. Uses of Ammonia. Ammonium hydroxide is used in the household for softening water for laundry purposes. We shall see later that this is brought about by causing the precipitation of the compounds present in the water which interfere with the action of soap. Ammonium hydroxide also acts as a cleansing agent by converting the oils and grease present into exceedingly fine glob- ules which do not adhere to one another and so can be removed by water. Ammonia is of value in cleaning certain metals from rusts and deposits caused by the presence of sulphur dioxide and other gases in the air; the use for this purpose is based on the fact that ammonia reacts with the salts present by direct addition and converts them into compounds soluble in water. Ammonia is taken internally as a medicine, and is a heart stimulant. It is the gas liberated by " smelling salts." In larger quantities it is an active poison. 348. Ammonium Salts. Ammonia unites directly with acids and forms ammonium salts; ammonium chloride is formed when ammonia and hydrogen chloride are brought into contact (140). This experiment was first performed by Priestley. He had iso- lated ammonia gas, which he called alkaline air, and hydrogen chloride, which he called acid air; and in an endeavor to get a neutral air he brought the two together and, as a result, discovered the composition of ammonium chloride, which was up to that time unknown. Ammonium salts are also formed by neutralizing ammonium hydroxide with solutions of acids, for example, ammonium sul- phate can be formed in this way: 2NH 4 OH + H 2 SO 4 = (NH 4 ) 2 SO 4 + 2H 2 O The salts derived from ammonia are called ammonium salts because they resemble in properties the salts of metals, the names of which AMMONIA AND ITS DERIVATIVES 315 in most cases terminate in ium, for example, sodium, calcium, aluminium. All ammonium salts contain the combination of atoms represented by the symbols NEU; this is called a radical and passes unchanged from one ammonium compound to another. It plays the same part in ammonium salts that metallic atoms play in their salts; it passes into the form of an ion when ammonium salts are dissolved in water, and has the valence 1. When attempts were made to isolate it, decomposition into ammonia and hydrogen took place. But some of these experiments lead to the conclusion that the radical possesses the properties of metals. When an electric current is passed through a solution of sodium chloride and the cathode is mercury, the metallic sodium liberated dissolves in the mercury and forms a compound with it, which is hard and shows the properties of an alloy. 1 When a solution of ammonium chloride is electrolyzed under the same conditions below a solid ammonium amalgam is obtained which resembles the corresponding product containing sodium. If the temperature is allowed to rise, the amalgam softens and swells up to a spongy mass as the result of the decomposition of the ammonium into ammonia and hydrogen. These facts indicate that the radical ammonium exhibits the properties of metals in alloys when it is in combination with another metal; and it has already been shown that the radical plays the part of a metal when it is in combination with acid radicals. A large number of ammonium salts are known and many of them have interesting applications. Ammonium chloride, sal ammoniac, is obtained in impure condition by neutralizing the gases given off when coal is distilled in making coke or illuminating gas. It is readily purified by sublimation as its vapor pressure is equal to the pressure of the atmosphere at about 338. Large quantities of the salt are used in the manufacture of dry cells, as a source of ammonia, and in chemical laboratories. It is used as a flux in soldering, as it serves to remove the coating of oxides formed on the solder and the metal and thus makes it possible for the two 1 Alloys are usually produced by melting two metals together; brass is prepared from copper and zinc in this way. Alloys often contain compounds of definite composition formed as the result of the union of the metals from which they are prepared. The alloys containing mercury are called amal- gams; sodium amalgam, for example, contains a compound of the formula NaHg a . 316 INORGANIC CHEMISTRY FOR COLLEGES to alloy. The ammonium chloride dissolves the oxides because when it is heated it dissociates into ammonia and hydrogen chloride, and the latter converts the oxides into chlorides, which melt and run off. Crude ammonium sulphate, (NH 4 )2SO 4 , is obtained from the gas works and is used chiefly as an ingredient of fertilizers. 349. During the recent war the supply of sulphuric acid in Germany was limited on account of the difficulty of obtaining sulphur and sulphides. Large quantities of ammonium sulphate were manufactured from calcium sulphate directly without first preparing sulphuric acid. Finely ground gypsum, the form in which calcium sulphate occurs in nature, was stirred in water and treated with ammonia and carbon dioxide. The ammonium carbonate produced reacted with the sulphate and formed am- monium sulphate and calcium carbonate: CaSO 4 + (NH 4 ) 2 CO 3 = (NH 4 ) 2 S0 4 + CaCO 3 After filtrating from the insoluble carbonate, the solution on evap- oration yielded ammonium sulphate. This process avoids the manufacture of sulphuric acid and is of economic significance. 350. Ammonium carbonate, (NH^COa, is formed by passing carbon dioxide through a solution of ammonia, if precautions are taken to have an excess of the latter present. On evaporation the normal carbonate is obtained, but it is quite unstable and when left in the air soon loses ammonia an 1 is changed into the bicar- bonate, NH 4 HC03. The latter can be made by completely saturating a solution of ammonia with carbon dioxide and subse- quent evaporation. The salt is obtained as white crystals, which slowly give off ammonia. This property leads to its use in " smell- ing salts." Ammonium bicarbonate is moistened with a solution containing a perfume of pleasant odor; when the bottle con- taining the salts is opened the ammonia which has accumulated is given off in the desired concentration. The uses to which ammo- nium sulphide and polysulphide (282, 311) and ammonium per- sulphate are put have been mentioned; other salts will be described later. 351. General Properties of Ammonium Salts. All these salts are decomposed when heated, but the behavior of any particular one is determined by the acid radical present. If the acid from AMMONIA AND ITS DERIVATIVES 317 which the salt is derived is volatile, the salt sublimes and if pure leaves no residue; if it is not volatile, the acid or its anhydride is left when the salt is ignited. It is for this reason that ammonium hydroxide is used frequently in analytical chemistry when it is necessary in the process to neutralize an acid ; the solution obtained can be evaporated to dryness and heated, and thus freed from the base used during the course of an analysis, in which the presence of sodium or potassium would interfere with obtaining the results desired. All ammonium salts with the exception of a few so-called double salts, like the one of the composition PtCL^NEUCl, are readily soluble in water. The fact that all ammonium salts are decomposed in aqueous solution by a base is utilized in the general test for these salts. The test is carried out by treating a solution of the material under examination with a solution of sodium hydrox- ide and noting the odor; if ammonia is not present the solution is heated and the odor again noted. All ammonium salts yield ammonia under these conditions. 352. Other compounds of Nitrogen and Hydrogen. It has been pointed out that nitric acid can be reduced by nascent hydrogen to ammonia. When the most active reducing agents are used with the acid or with the oxides of nitrogen, the final reduction-product in all cases is ammonia. When the less active reducing agents are used, however, under the proper conditions other reduction-products are formed. For example, when nitric oxide, NO, is reduced by combining the action of potassium sul- phite with that of the nascent hydrogen generated by the action of water on sodium amalgam (note, 348), a compound of the formula N2H4, called hydrazine y is formed. In this compound the reduction of the nitrogen has not proceeded as far as is the case with ammo- nia; two nitrogen atoms in the form of ammonia are in combi- nation with six hydrogen atoms, whereas in hydrazine two nitrogen atoms are united with four atoms only. If hydrazine is treated with hydrogen generated by a metal and an acid it is reduced further to ammonia. Hydrazine is a colorless liquid which boils at 113. It unites with water to form a compound which unlike ammonium hydrox- ide is stable. Hydrazine and its compound with water resemble ammonia and ammonium hydroxide in chemical properties. 318 INORGANIC CHEMISTRY FOR COLLEGES Since it contains two atoms of nitrogen in the molecule hydrazine can form salts by uniting with either one or two molecules of a monobasic acid. Hydrazine is readily oxidized and is, therefore, an active reducing agent. 353. When dilute nitric acid is reduced at low temperatures with nascent hydrogen generated by the action of tin on the acid, all but one of the oxygen atoms are removed and three hydrogen atoms are added. The resulting product is called hydroxylamine, and has the formula NH 2 OH. It is a white solid which melts at 33. The compound forms salts by direct addition with acids, and its aqueous solution shows the properties of a weak base. Like hydrazine, hydroxylamine is an active reducing agent. 354. Hydrazoic acid, HN 3 , is a compound of peculiar interest on account of the fact that it is so unstable. It can be formed by the oxidation of hydrazine with n;trous acid: N 2 H 4 + HNO 2 = HN 3 + 2H 2 O The sodium salt of the acid is more conveniently prepared by treating sodium amide with nitrous oxide: NaNH 2 + N 2 O = NaN 3 + H 2 O Since the acid is volatile with steam it can be obtained by distilling the sodium salt with dilute sulphuric acid. Hydrazoic acid, which is also called hydronitric acid, is a colorless liquid which possesses a strong, disagreeable odor; it boils at 37 and explodes with great violence when heated. The reaction is accompanied by the evo- lution of a large quantity of heat: 2HN 3 = 3N 2 +H 2 + 2 X 61,600 cals. The salts of the acid, which are called azides, are, in the case of the heavy metals, explosive compounds. Lead azide has recently come into use in the manufacture of caps for high explosive shell. EXERCISES 1. Ammonia was passed over hot copper oxide and the nitrogen formed was determined: 2NH 3 +3CuO= N 2 -f 3H 2 O + 3Cu. (a) Calculate the percentage composition of NH 3 from the following results: Weight CuO before experiment 6.450 grams; weight CuO after experiment 6.258; volume of N 2 at and 760 mm. 89.6 c.c. (6) Calculate the percentage composition of NH 3 from its formula, AMMONIA AND ITS DERIVATIVES 319 2. The substances having the following formulas are commonly used in the laboratory as drying agents: CaCl 2 , KOH, CuSO 4 , H 2 SO 4 , P 2 O 6 . Which can and which cannot be used to dry ammonia? Give a reason in each case. 3. (a) Calculate from its formula the weight of 1 liter of NH 3 . (6) Cal- culate the specific gravity of NH 3 compared with air, 1 liter of which weighs 1.298 grams, (c) Could NH 3 be used in balloons? What evident disadvan- tage would the gas possess for this purpose? 4. How could you determine the volume of NH 3 in a mixture of the latter and air contained in a flask? 5. Write equations for three reactions by which NH 3 may be produced from three different salts, using a different base in each case. 6. Compare the action at high temperatures of H^O and NH 3 on metals. 7. State three ways in which you could show that a solution of NH 3 in water contains OH ions. 8. What would you expect to be formed when an electric current is passed through a solution of (NH 4 )2SO 4 in water? 9. How could you distinguish from one another the following: (a) (NH 4 ) 2 SO 4 , (6) NH 4 C1, (c) NH 4 NO 3 , (d) NaCl, (e) Na 2 SO 4 , (/) NaOH, (0) NH 4 HC0 3 , (A) (NH 4 ) 2 S? 10. Name the compounds and ions present in a solution of NH 3 in water and represent by equations the equilibrium that exists between them. Show how the equilibrium changes in each reaction when the solution is heated to boiling, CHAPTER XXIII NITRIC ACID, NITROUS ACID, AND THE OXIDES OF NITROGEN 355. Nitric acid, nitrous acid and the oxides of nitrogen are of particular interest and significance; their chemical properties illustrate some of the most fundamental principles underlying the behavior of molecules, and the energy changes involved in their transformations furnish an opportunity to grasp clearly the sig- nificance of the energy factor in chemical change. We have studied in some detail carbon monoxide, CO, and carbon dioxide, CO2, and have learned that these compounds can be readily formed, and that in their production large quantities of heat are evolved the reactions are exothermic. It will be of interest to study now the nitrogen compounds of analogous composition, nitric oxide, NO, and nitrogen dioxide, NO2, because these compounds are endo- thermic that is, heat is absorbed when they are formed. The tendency in nature is to render energy unavailable a stone falls to the ground and the power it had to do work due to its position is lost, and the energy passes into heat, only a part of which can be utilized. Reactions which continue to take place when once started are, in most cases, those which give off energy. Carbon is a storehouse of chemical energy that can be utilized to do the mechanical work of the world; and the product formed when it burns freely is comparatively inert; millions of tons of carbon dioxide the product of burning coal are allowed to escape, because the compound contains such a small amount of available energy. Nitrogen, on the other hand, unites with oxygen only as long as we continue to supply a large amount of energy; and the resulting oxide can be used as a source of energy. It is an endothermic compound, and as it is an excellent example of the class to which it belongs its study will be of great importance. The tremendous, destructive force shown by gunpowder, dynam- ite, smokeless powder, and the other high explosives used in warfare 320 NITRIC ACID, NITROUS ACID, OXIDES OF NITROGEN 321 and for blasting and other industrial purposes, is traceable to the fact that when a nitrogen atom unites with oxygen atoms the energy furnished to effect the union is stored up and becomes available when the substance containing these atoms reacts with other compounds. Since in all chemical changes matter and energy are involved, it is important not to lose sight of the energy factor when considering the matter, which, in most cases, more often affects our senses. In all industrial problems involving the prep- aration of chemical compounds, the skilled chemical engineer pays particular attention to the energy involved, since it is often an important factor in the cost of the manufactured product. NITRIC ACID 356. Nitric acid, HNOs, does not occur in nature in the free condition except in traces in the air after a thunderstorm (324). It is active chemically and if set free would soon find something with which to react. It occurs as nitrates in certain parts of the earth. Potassium nitrate is found in the soil near cities in certain Oriental countries, having been produced as the result of the action of nitrifying bacteria on animal refuse. Such soil was the source of Bengal saltpeter. Large deposits of sodium nitrate are found in Chile, and these are now the chief source of the combined nitrogen used in the world. The deposits are in an arid region of North Chile and cover about 500 square miles; they are from 2 to 10 feet thick and contain up to 60 per cent of sodium nitrate. Guano (315), formerly much used as a fertilizer, contains a large percentage of nitrates which have been produced as the result of bacterial action on the organic nitrogen compounds present in the excreta of birds, from which the guano was formed. Nitrates which are present in the soil are used by the plant in building up the organic nitrogen compounds produced in their growth. If the material grown is taken away, in a short time the soil becomes depleted and it is necessary to renew the nitrates through the application of fertilizers. Nitrates are found in small and varying amounts in water, and their presence and quantity serve, as we shall see, to indicate to the chemists the history of the water as regards pollution. Nitric acid was known in ancient times, and was called aqua 322 INORGANIC CHEMISTRY FOR COLLEGES fortis on account of the fact that it was an excellent solvent for substances which were not affected by less powerful reagents. Lavoisier (1776) first showed that it contained hydrogen, nitrogen, and oxygen, and Gay-Lussac later determined its exact composi- tion. 357. Preparation of Nitric Acid. In preparing nitric acid from its salts we make use of the important fact that double decompositions proceed to completion provided one of the products of the reaction is a gas. Nitric acid is a liquid at the ordinary temperature but boils at 86; as a consequence, if a nitrate is mixed with sulphuric acid and heated to this temperature the conditions are those which lead to double decomposition. The equation for the reaction used in the technical preparation of nitric acid is as follows : NaNO 3 + H 2 SO 4 = NaHSO 4 + HNO 3 The proportions indicated are used rather than those which lead to the formation of neutral sodium sulphate: 2NaNO 3 + H 2 SO 4 = Na 2 SO 4 + 2HNO 3 This second reaction appears to be the more economical one because 1 molecule of sulphuric acid causes the liberation of 2 molecules of nitric acid, whereas in the first reaction but half that amount would be obtained; the cost of the sulphuric acid used in the process based on the second equation would be, thus, one-half the cost of the acid used in the first one. It will be recalled that when a reaction takes place according to the second equation it proceeds in steps; that represented by the first equation first takes place, and then the acid sulphate reacts with another molecule of nitrate as follows: NaNO 3 + NaHSO 4 = Na 2 SO 4 + HNO 3 The temperature at which this double decomposition takes place is much higher than that required for the first step; and as this temperature is above that at which nitric acid partially decomposes the reaction cannot be used. In manufacturing nitric acid, sodium nitrate is heated in cast-iron retorts with sulphuric acid having a specific gravity of 1.71 to 1.83. depending on the strength of acid desired. The NITRIC ACID, NITROUS ACID, OXIDES OF NITROGEN 323 escaping vapor is condensed in glass, stone-ware, or aluminium pipes surrounded by water. To free the acid from the small amounts of oxides of nitrogen formed in the distillation, which give it a yellow color, air is blown through the acid. If concentrated nitric acid and concentrated sulphuric acid are mixed and heated cautiously, nitric acid free from water distills over. It contains nitrogen dioxide produced as the result of the decomposition of a part of the nitric acid and, as a result, has a brown color. The acid fumes in the air and is called, therefore, fuming nitric acid ; it has the specific gravity 1 .6 and is a powerful oxidizing agent. 358. Synthetic Nitric Acid. The fact has been known for a long time that when an electric spark is passed through a mixture of nitrogen and oxygen the elements unite and, if water is present, nitric acid is formed. Although Cavendish carried out this experiment in 1785 it has only been in recent years that it has been possible to apply the reactions involved to the manufacture of nitric acid. Science had to advance as the result of the discovery of the fundamental laws underlying chemical equilibrium, before it was possible to develop a great chemical industry out of a change of theoretical interest only. The formation of nitric acid from air and water is based on reactions represented by the following equations : N 2 + 2 + 43,200 cals *=* 2NO 4NO + 3O 2 + 2H 2 O = 4HNO 3 The second reaction takes place readily; if nitric oxide and an excess of air are passed through water at ordinary temperatures nitric acid is formed. The first reaction is the one that required extensive study before it could be utilized for large-scale production. The reaction is a reversible one, and, consequently, to discover the most favorable conditions for its use the equilibria at different temperatures had to be determined, and the rate of the reaction at these temperatures, investigated. Since the formation of nitric oxide is an endothermic reaction, rise in temperature causes the equilibrium to shift in such a way that the proportion of nitric oxide in the resulting gases increases (law of mobile equilibrium). With rise in temperature the rate at which equilibrium is attained increases, 324 INORGANIC CHEMISTRY FOR COLLEGES Electrode for Arc...,, The study of the equilibrium from these two points of view disclosed the best practical conditions for carrying out the reaction. It was found, for example, that at 1538 the gases in equilibrium when air was used contained 0.37 per cent nitric oxide, at 1922, 0.97 per cent, and at 2927, 5.0 per cent. The rate at which equilibrium was attained at the high temperatures required was measurable but high enough for practical purposes, and it dropped off rapidly with falling temperature. The data obtained led to the conclusion that the reaction should be carried out at the highest temperature possible, that no catalytic agent was neces- sary, and that after the gases had reached equilibrium they should be suddenly cooled in order to prevent the shifting of the equilibrium, which would result, with slowly falling temperature, in the decomposition of the nitric oxide produced. A number of forms of apparatus have been devised to meet these conditions. Since the temperature required is so high, an electric arc is the source of heat in all of them. In the Birkeland-Eyde process used in Norway, an arc between two carbon poles is flattened out into a fan-like discharge by means of an electro-mag- net. A cross-section of the apparatus is shown in Fig. 31. The gases from the furnace, after cooling, are passed through towers filled with tiles down which water trickles. Arrin/ef- '^Sas and Air Outlet The dilute nitric acid obtained FlG - 31. is neutralized with lime, CaO, and is converted into calcium nitrate, which is obtained on evaporation of the solution. The yield is said to be 70 grams of nitric acid per kilowatt hour. It is of interest to compare here the process employed to make nitrogen and hydrogen unite to form ammonia (338) with that used to prepare nitric oxide from nitrogen and oxygen. In the first case the reaction is exothermic and in the second case it is endothermic; and this difference led to quite different procedures NITRIC ACID, NITROUS ACID, OXIDES OF NITROGEN 325 in using the reactions for the two syntheses. The nitrogen- hydrogen reaction was carried out at as low a temperature as pos- sible, and as its rate decreased as the temperature was lowered it was necessary to use a catalyst. The nitrogen-oxygen reaction was carried out at as high a temperature as possible, and as increase in temperature is associated with increase in rate, no catalyst was required. In the nitrogen-hydrogen reaction there is a change in volume as the result of the union 4 volumes become 2 and, as a consequence, the synthesis of ammonia was carried out at the highest practical pressure. In the nitrogen-oxygen reaction there is no change in volume and the synthesis was carried out at ordinary pressures. In the synthesis of nitric oxide energy is absorbed and the cost of electrical power is a factor in the cost of the nitric acid manufactured. In the synthesis of ammonia energy is given off; the power necessary is only that required to compress the gases, and we have seen that a part of this is recovered. It is important to emphasize the fact that it was the application of the principles underlying mobile equilibrium that made it possible to develop into commercial processes reactions which under ordinary circumstances furnished but an exceedingly small frac- tion of a per cent of the desired product. 359. The most recently developed method of manufacturing nitric acid is based on the oxidation of ammonia. The economic significance of this process has already been mentioned and it was pointed out that it was used during the recent war. The process is simple in principle. When a mixture of ammonia and air is passed over a proper catalyzer, nitric oxide is formed as the result of a reaction between the ammonia and the oxygen in the air: 4NH 3 + 5O 2 = 4NO + 6H 2 O When the nitric oxide is cooled it reacts with more oxygen and nitrogen dioxide, NO2, is formed; the latter, when mixed with air and passed into water, is converted into nitric acid : 4NO 2 + 2H 2 O + O 2 = 4HNO 3 The oxidation of ammonia is carried out in chambers lined with fire-brick in which several layers of the catalyzer are placed. In some plants fine platinum gauze is used as the contact agent and in others an oxide of iron. The reaction goes to completion at the 326 INORGANIC CHEMISTRY FOR COLLEGES temperature used, about 600, and being exothermic does not require the application of external heat after it has been started. The nitric oxide is cooled, mixed with more air, and passed through towers lined with tile down which water trickles. In order to effect complete absorption of the gases several towers are used, the last containing a solution of sodium carbonate. Several hundred thousand tons of nitric acid and nitrates have been made in a single year by this process. . ; ' 360. Physical Properties of Nitric Acid. Nitric acid is a color- less liquid which boils at 86, and has the specific gravity 1.56 at 0. It usually has a yellow color which is produced by the pres- ence of oxides of nitrogen formed as the result of the decomposition of some of the acid. The concentrated nitric acid of commerce is the constant boiling mixture of the acid and water; it contains 68.6 per cent nitric acid, has the specific gravity 1.41 at 15 and boils at 120.5. 361. Chemical Behavior of Nitric Acid. The acid decomposes slowly at its boiling-point according to the following equation: 4HNO 3 = 2H 2 O -f 4NO 2 + O 2 Nitric acid is miscible with water in all proportions and the solutions show strong acidic properties. The acid is dissociated into H + and NO 3 ~ ions, and, as is usual, the extent to which the dissociation takes place is determined by the concentration of the acid; in one-tenth normal solution at 18 it is 92 per cent disso- ciated; the ionization of hydrochloric acid at this dilution is 92 and of sulphuric acid 61 per cent. The ionization in a solution containing 62 per cent of nitric acid is only 9 per cent. The extent to which ionization has taken place in a solution of nitric acid is an important factor in its chemical behavior; undissociated nitric acid is an active oxidizing agent, whereas the ions produced from it, NOa~, do not show this property. Anhydrous nitric acid is such a powerful oxidizing agent that a flame is sometimes produced when it is brought into contact with oxidizable substances. This can be readily shown by simple experiments. If warm nitric acid is poured on sawdust which has been slightly heated, the oxi- dation is so violent that the wood ignites and burns with a flame. The dense brown fumes produced are nitrogen dioxide formed from NITRIC ACID, NITROUS ACID, OXIDES OF NITROGEN 327 the nitric acid. If a bit of wool is placed in the mouth of a test- tube containing some pure nitric acid, and the latter is then heated, the vapors given off from the acid ignite the wool, which burns with a flame. It is evident that great care must be used in ship- ping and storing pure nitric acid. The danger of fire from con- centrated nitric acid is not so great, but care is necessary with it, nevertheless, because it is highly corrosive. 362. The reduction-product formed from nitric acid when it oxidizes is, as we would conclude, dependent on the strength of the acid and the activity of the reducing agent with which it interacts. In most cases, however, the chief reduction-product is nitric oxide, the acid breaking down as indicated by the following equation: 2HNO 3 = 2NO + H 2 O + 3O It must be emphasized that this is not an equation expressing the decomposition of nitric acid alone; it represents the decomposition which takes place when nitric acid is brought into contact with some oxidizable substance, and it serves, therefore, as a partial equation in a series representing the oxidation of a substance by nitric acid. We shall see that the equation is used repeatedly when oxidation reactions are written. 363. Nitric acid oxidizes all the metals except the so-called noble metals, gold and platinum, for example. The nitrates formed are soluble in water, and, consequently, nitric acid dis- solves these metals. It was on account of this solvent action of the acid that it was called by the alchemists " aqua fortis." In writing equations for the reactions involved, it is best to separate them into steps and combine the partial equations in the way already explained (287, 288). For example, the oxidation of cop- per by nitric acid can be represented as follows ; 2HNO 3 = 2NO + H 2 O + [3O] 3Cu + [30] = [3CuO] [3CuO] + 6HNO 3 = 3Cu(N0 3 ) 2 + 3H 2 3Cu + 8HNO 3 = 3Cu(NO 3 ) 2 +.4H 2 O + 2NO 328 INORGANIC CHEMISTRY FOR COLLEGES The oxidation of silver, which has the valence 1, can be written in a similar way: 2HNO 3 = 2NO + H 2 O + [3O] 6Ag + [30] = [3Ag 2 0] [3Ag 2 O] + 6HNO 3 = 6AgNO 3 + 3H 2 O 6Ag + 8HNO 3 = 6AgN0 3 + 4H 2 O + 2NO In this case all the integers indicating the number of molecules can be divided by two, and the equation is simplified to read as follows : 3Ag + 4HNO 3 = 3AgNO 3 + 2H 2 O + NO Many acid-forming elements are oxidized by strong nitric acid; the equation for the reaction with sulphur is as follows : 2HN0 3 = 2NO + [H 2 0] + [3O] S + [30] = [S0 3 ] [S0 3 ] + [H 2 0] = H 2 S0 4 S + 2HNO 3 = 2NO + H 2 SO 4 Carbon when heated with strong nitric acid is oxidized: 4HNO 3 = 4NO + 2H 2 O + [6O] 3C + [6O] = 3CO 2 3C + 4HNO 3 = 4NO + 2H 2 O + 3CO 2 The nitric oxide formed in all the reactions mentioned is converted into nitrogen dioxide if air is present, because the colorless nitric oxide unites at ordinary temperatures with the oxygen of the air to form nitrogen dioxide, which is a brown gas. If the oxidations are carried out in the presence of an excess of concentrated nitric acid, the dioxide is also formed. When the most active metals react with nitric acid, hydrogen is formed; if the solution is very dilute and the acid almost com- pletely in the form of ions, the hydrogen escapes; in slightly stronger solutions the hydrogen is oxidized by the nitric acid to water and the acid is reduced to ammonia, which forms ammonium nitrate with some of the acid present. Under special conditions nitric acid when it acts as an oxidizing agent may be reduced to NITRIC ACID, NITROUS ACID, OXIDES OF NITROGEN 329 nitrous oxide, hydroxylamine, or hydrazine (352) ; the reaction to be definitely remembered, however, is its reduction to nitric oxide. Many compounds are oxidized by the acid; the reaction in the case of sulphur dioxide is expressed by the following equations : 2HNO 3 = H 2 O + 2NO + [3O] 3S0 2 + [30] = [3S0 3 ] [3S0 3 ] + 3H 2 = 3H 2 S0 4 2HN0 3 + 3SO 2 + 2H 2 O = 2NO + 3H 2 SO 4 When nitric acid oxidizes hydrochloric acid, the nitric oxide formed unites with some of the chlorine produced by the oxidation to form nitrosyl chloride, NOC1: 2HN0 3 = H 2 O + [2NO] + [3O] 6HC1 + [3O] = 3H 2 O + 6C1 [2NO] + 2C1 - 2NOC1 2HNO 3 + 6HC1 - 4H 2 O + 2C1 2 + 2NOC1 HNO 3 + 3HC1 = 2H 2 O + C1 2 + NOC1 It will be seen from the equations that chlorine is formed. Since chlorine reacts with many substances and converts them into chlorides which are soluble, the mixture of hydrochloric acid and nitric acid is an excellent solvent. The mixture was called aqua regia by the alchemists because it dissolved the royal metal, gold. Aqua regia is much used in the chemical laboratory to convert various insoluble substances into soluble compounds. It is gen- erally prepared by mixing 1 volume of concentrated nitric acid and 3 volumes of concentrated hydrochloric acid. 364. A solution of nitric acid in water shows all the charac- teristic properties of an acid; it reacts with oxides and hydroxides of metals to form nitrates. There are certain organic compounds called alcohols which contain the hydroxyl group and interact with strong acids to form salts. Common grain alcohol has the formula C 2 H5OH it is the hydroxide of the radical C 2 H5, which is called ethyl. Alcohol reacts with nitric acid and forms ethyl nitrate : . C 2 H 5 OH + HN0 3 = C 2 H 5 N0 3 + H 2 O 330 INORGANIC CHEMISTRY FOR COLLEGES Since alcohol can be oxidized by nitric acid, special precautions are necessary in carrying out the reaction to prevent the oxidation which often takes place with explosive violence. Glycerine is an alcohol of the formula C 3 H5(OH) 3 and when treated with a mixture of nitric and sulphuric acids forms a nitrate, C 3 H5(NO 3 ) 3 , which is commonly called nitroglycerine. The product, which is a liquid, is made in large quantities and, mixed with sawdust, clay, and other substances is used under the name of dynamite as an ex- plosive. Cellulose, which is the chief constituent of raw cotton and wood, is also an alcohol. It forms a nitrate with nitric acid : C 6 H 7 02(OH)3 + 3HNO 3 = C 6 H 7 O 2 (NO 3 ) 3 + 3H 2 O The nitrate is commonly called nitrocellulose and is the chief ingredient of smokeless powder. Nitric acid reacts with many hydrocarbons and their deriva- tives and forms the so-called nitro compounds, all of which contain the characteristic nitro group, N0 2 . For example, it reacts with benzene as follows: C 6 H 6 + HNO 3 = C 6 H 5 NO 2 + H 2 O The product, nitrobenzene, can be reduced by nascent hydrogen to aniline, CeHsNH^, which is used in the manufacture of the so-called aniline dyes. With toluene, a derivative which contains three nitro groups can be made: CH 3 .C 6 H 5 + 3HN0 3 - CH 3 .C 6 H 2 (N0 2 ) 3 + 3H 2 O Trinitrotoluene, commonly called T.N.T., is a very important high explosive used in shells for military purposes. Picric acid, which is a valuable explosive, is also used in large quantities in warfare. It is trinitrophenol, C6H 2 (OH)(NO 2 ) 3 , and is made from phenol, CeBkOH, which is often called carbolic acid. 365. When all the explosives which have been mentioned de- compose, there is a liberation of the energy which was stored up in the compound as the result of the union of nitrogen and oxygen in the nitric acid from which they were prepared ; the carbon present is oxidized to carbon dioxide and carbon monoxide, or if there is not enough oxygen present in the molecule to oxidize all the carbon, NITRIC ACID, NITROUS ACID, OXIDES OF NITROGEN 331 it is liberated as such and appears as black smoke. Some explo- sives, like nitroglycerine, decompose very readily; others can be handled with safety, since it is necessary to detonate them before they explode. A detonater is a compound that explodes when it is subjected to a shock such as the blow of the hammer of a rifle. The decomposition which takes place sets up vibrations which cause the explosion of compounds not readily influenced by simple mechanical shocks. 366. Hair, flesh, feathers, silk, and other nitrogenous constit- uents of living things, which are called proteins, are formed as the result of the union of a large number of complex molecules one of which is derived from benzene. When any of these substances is treated with strong nitric acid, a nitro compound is formed just as benzene yields nitrobenzene. The compound, like most nitro compounds, is yellow. It is for this reason that nitric acid pro- duces a yellow stain when left in contact with the skin or woolen fabrics. When acids are allowed to stay in contact with dyed cloth a red stain generally develops as a result of a chemical change in the dye. Such a stain can be removed by treating it with ammonia; the dye in this case acts as an indicator and the effect of an acid is neutralized by the application of a base. Ammonia is used for this purpose because it does not materially affect the fabric, and it is volatile and any excess is removed by evaporation. Sodium hydroxide decomposes proteins and, therefore, attacks wool and silk. When a stain produced by nitric acid is treated with ammonia the original color is not restored because the yellow nitro compound produced by the acid is not destroyed. 367. Nitrates. All normal nitrates are soluble in water, and the difficultly soluble basic nitrates are soluble in the presence of an excess of nitric acid. All nitrates decompose when heated, those of the more active elements requiring the higher temperatures to effect decomposition. In the case of sodium and potassium, the nitrates can be melted without decomposition, but at red heat they lose oxygen and pass into nitrites; sodium nitrite may be formed in this way: 2NaNO 3 = 2NaNO 2 + 2 The nitrates of all metals except those that resemble sodium in chemical properties break down into the oxide of the metal and 332 INORGANIC CHEMISTRY FOR COLLEGES the anhydride of the acid. The anhydride of nitric acid is NoOs: 2HNO 3 = N 2 O 5 + H 2 O Nitric acid anhydride is unstable at the temperature at which nitrates decompose and breaks down into nitrogen dioxide and oxygen: 2N 2 O 5 - 4NO 2 + O 2 The equation which represents the decomposition of copper nitrate when heated is, therefore: 2Cu(NO 3 ) 2 = 2CuO + 4NO 2 + O 2 368. The test for nitrates is based on the production of a color, and not, as usual, on the formation of some insoluble compound, as is the case with sulphates, chlorides, etc. The formation of a characteristic insoluble compound cannot be used in this case because of the fact that nitrates are soluble. We have learned that when nitric acid oxidizes substances it is reduced to nitric oxide. This oxide forms a compound with ferrous sulphate, FeSO4, which is highly colored in great dilution, and its formation, therefore, serves as a delicate test for nitric oxide and for nitric acid provided the latter is brought into contact with something that reduces it to nitric oxide. Ferrous sulphate is oxidized by nitric acid to ferric sulphate, so it serves two purposes in the test based on these reactions. The test is carried out as fol- lows: A solution of the substance to be tested is mixed in a test-tube with a strong solution of ferrous sulphate. The tube is inclined and concentrated sulphuric acid cautiously poured in. As the acid flows down the side of the tube it sinks to the bottom and forms a heavy layer beneath the aqueous solution. If a nitrate is present a brown ring appears at the juncture of the two liquids. The sulphuric acid liberates nitric acid from the nitrate; the acid oxidizes a part of the ferrous sulphate and is thereby reduced to nitric oxide, and finally the nitric oxide unites with some of the ferrous sulphate to form a compound of the formula FeSCU NO, which possesses a brown color. The test is a delicate one because it shows the presence of a very small amount of nitric acid or a nitrate. The individual salts of nitric acid will be described in connec- NITRIC ACID, NITROUS ACID, OXIDES OF NITROGEN 333 tion with the discussion of the chemistry of the metals which they contain. Ammonium nitrate was used in large quantities as a high explosive during the recent war. The compound under ordinary circumstances is inert. When the dry salt is heated it decomposes into nitrous oxide and water, NH 4 NO 3 = N 2 O + 2H 2 O, and the reaction serves as the most convenient method of pre- paring the gas. When, however, the salt is detonated it decom- poses with explosive violence. One of the most valuable fillings for high explosive shells is a mixture of ammonium nitrate and T.N.T.; it is called amatol. NITROUS ACID 369. Nitrous acid is used in making the so-called azo dyes, and as these are readily prepared and furnish a great variety of colors and shades, large amounts of the salts of the acid are manufactured. The acid itself is very unstable, as it breaks down largely into its anhydride and water when liberated. The chemical relation between nitrous acid, HNC>2, and nitric acid, HNOs, is similar to that between sulphurous acid, H^SOs, and sulphuric acid, H2SO4, the similarity being traceable to the fact that in the two cases the acids differ in composition by one oxygen atom. We shall see later that there are many cases in which one element forms two acids containing different amounts of oxygen, and that the differ- ence in their chemical properties can be traced to this cause. It is well to emphasize the relationship here. 370. The methods commonly used to prepare nitrous acid and nitrites are not analogous to those used in the case of the sulphur compounds for a number of reasons. The supply of raw materials is different in the two cases, and as the reactions are endothermic in one case and exothermic in the other, the ease of preparation of the nitrogen compounds and the sulphur compounds from the elements is markedly different. It will be recalled that sulphur is the raw material from which sulphurous acid and sulphuric acid are made, and that it is only recently that nitrogen has been used as the source of nitric acid. Nitrous acid and its salts are now manufactured from the oxides produced by burning nitrogen, 334 INORGANIC CHEMISTRY FOR COLLEGES but the older method starting with the sodium nitrate that occurs ready made in nature is the one commonly used. When sodium or potassium nitrate is heated to a high tempera- ture a part of the oxygen is lost, and a nitrite is formed : 2NaNO 3 = 2NaNO 2 + O 2 In manufacturing nitrites lead is usually stirred with the molten mass; it helps in the reduction by uniting with the oxygen, and is converted into lead oxide, PbO. When the reaction is complete the mixture is allowed to cool and is treated with hot water; the solution on evaporation yields sodium nitrite in the form of colorless crystals. 371. When a solution of sodium nitrite is treated with an acid a reaction takes place which is analogous to that in the case of sodium sulphite, and the reasons underlying the double decompo- sition are the same in the two cases (287) . Nitrous acid is formed and in part decomposes into nitrous anhydride and water; 2NaNO 2 + H 2 SO 4 = Na 2 SO 4 + 2HNO 2 2HNO 2 <=> N 2 O 3 + H 2 O The nitrous anhydride escapes as a brown gas. When nitrous anhydride is passed into a solution of sodium hydroxide it reacts and forms a nitrite, just as sulphur dioxide under the same conditions forms a sulphite : H 2 O + N 2 O 3 <=* 2HNO 2 NaOH + HNO 2 = NaNO 2 + H 2 O Like sulphurous acid, nitrous acid is a mild oxidizing and reducing agent. The reaction which takes place between hydriodic acid and sulphurous acid (sulphur dioxide and water) has already been given : 4HI + SO 2 = 2I 2 + 2H 2 O + S The reaction between hydriodic acid and nitrous acid (nitrous anhydride and water) is similar to the one just given, although nitrogen is not formed in this case: 2HI + N 2 3 = I 2 + H 2 + 2NO NITRIC ACID, NITROUS ACID, OXIDES OF NITROGEN 335 The oxidation of hydriodic acid can be traced to the same cause in the two cases. Both nitrogen and sulphur can exist in com- pounds in which the elements have lower valencies than in sulphurous and nitrous anhydrides. In sulphur dioxide and sul- phurous acid sulphur has the valence 4, and when these com- pounds serve as oxidizing agents the valence of the sulphur falls to zero when free sulphur is formed. When nitrous acid or its anhydride act as oxidizing agents they do not give up all the oxygen with which the nitrogen is combined, but are reduced to nitric oxide, the change in valence being from 3 to 2. 372. When sulphur dioxide acts as a reducing agent it unites with more oxygen and sulphur trioxide (sulphuric acid) is formed, the valence change being from 4 to 6 : SO2 > SOs. When nitrous anhydride serves as a reducing agent it passes to nitric anhydride (nitric acid) and the valence change is from 3 to 5: N2Oa ^Os. Sulphurous and nitrous acids act as reducing agents only when brought into contact with powerful oxidizing agents. When an acidified solution of potassium permanganate is treated with either substance the purple color of the salt disappears; the perman- ganate is reduced to a colorless compound and the sulphurous acid is oxidized to sulphuric acid or the nitrous acid to nitric acid. While there is a striking analogy between the nitrogen and sulphur compounds, there is a difference between them in their activity as reducing or oxidizing agents; sulphurous anhydride is a more active reducing agent than nitrous anhydride, and nitric acid a more active oxidizing agent than sulphuric acid. 373. Nitrites. The nitrites of most of the metals are known; they are all soluble in water, and yield nitrous anhydride when treated with an acid. The formation of this gas is used in the test for nitrites. The solution is acidified and warmed if necessary; if a nitrite is present' a yellow-brown gas is evolved. When the amount of nitrite present is small and the result indefinite on account of the small amount of gas given off, the liquid or the vapor from it is tested with a bit of paper which has been moistened with a solution containing potassium iodide and starch. The nitrous anhydride oxidizes the hydriodic acid and sets free iodine from the iodide, according to the reaction which has already been explained (371) , and the liberated iodine produces with the starch a charac- teristic blue color. The presence of the starch makes the test 336 INORGANIC CHEMISTRY FOR COLLEGES more delicate, for an amount of iodine that could not be recognized by its color produces a marked blue with starch. Nitrites, as might be expected, are much more stable under the influence of heat than nitrates. The nitrites of potassium and sodium resist very high temperatures; those of the heavy metals decompose at red heat or lower and yield oxides of the metals and nitric oxide and oxygen, which are produced as the result of the de- composition of nitrous anhydride. 374. Water Analysis. Having acquired a knowledge of the chemistry of ammonia, nitrous acid, and nitric acid we are now in a position to understand the significance of the methods used in the chemical analysis of water for sanitary purposes. One of the chief sources of contamination of water is sewage, which consists largely of animal refuse matter. A study of the amount of this present in a sample of water and the changes it has undergone gives a good indication of the purity of the water and its value for household use. The complex organic nitrogen compounds undergo decom- position slowly as the result of oxidation which is brought about through the influence of certain bacteria always present. This process of nitrification takes place in steps; first ammonia is set free by one kind of nitrifying bacteria, then this is oxidized to nitrous acid by another kind, and, finally, the nitrous acid is con- verted into nitric acid by a third kind. As these processes occur slowly it is possible, by determining quantitatively the amount of undecomposed material, ammonia, nitrous acid, and nitric acid present in a sample of water, to estimate the extent to which con- tamination has taken place, and whether it has been recent or not. By the time all the nitrogen has been converted into nitrates the impurities have disappeared and the water may be considered safe. In analyzing a sample of water a measured quantity of it is first distilled, and the amount of the free ammonia which passes over with the steam is determined. Potassium permanganate, KMnO4, is next added and the distillation continued. Potassium permanganate oxidizes the organic matter, and the nitrogen present is liberated as ammonia, the amount of which is determined as before. This ammonia is a measure of the undecomposed organic matter present in the water and is called albuminoid ammonia, on account of the fact that it is derived from the so-called albumens present. Nitrites and nitrates are determined in separate samples of NITRIC ACID, NITROUS ACID, OXIDES OF NITROGEN 337 the water under examination. All the quantitative determinations are made with the aid of substances that produce colors when brought into contact with the compound tested for. Colorimetric methods are necessary in the analysis because such small quanti- ties of the materials to be determined are present; we can see a color distinctly when it would be impossible to weigh the minute amount of material producing the color. The amount of the nitrogen compound present is determined by the depth of color produced. In each case the solution is placed in a tube and the intensity of the color compared with a series of standards made up from samples of water to which have been added known quan- tities of nitrites or nitrates. Below are given the results of the analysis of a sample of the water supplied to New York City. ANALYSIS OF CROTON WATER Parts per million Free Ammonia . 015 Albuminoid ammonia . 170 Nitrogen in nitrites . 000 Nitrogen in nitrates . 250 Chlorine in chlorides 2 . 100 Mineral matter 66.000 Total solids 81 .000 The chlorine which is present in a water chiefly as sodium chloride is determined, as it is also an indication of contamination if exces- sive amounts are found. OXIDES OF NITROGEN 375. There are five oxides of nitrogen which form a complete series in which nitrogen shows the valence of 1 to 5 inclusive; their formulas are N 2 O, NO, N 2 O 3 , N0 2 , and N 2 O 5 . In addition, there is an oxide of the formula N 2 O4 which is a polymer of nitrogen dioxide, NO 2 . 376. Nitrous Oxide. All of the oxides of nitrogen can be prepared from nitric acid. With the exception of nitrogen pen- toxide, which is obtained by the elimination of water from the acid, the compounds are made by reduction. By varying the concen- tration of the acid by dilution with water, and, therefore, its oxi- dizing power, and by selecting reducing agents of increasing activity 338 INORGANIC CHEMISTRY FOR COLLEGES we can prepare the oxides in which nitrogen has the valence 4, 3, 2, or 1. This method is not the most convenient one in all cases, however. Nitrous oxide is formed along with other oxides when dilute nitric acid is treated with zinc. Priestley, who discovered the gas in 1772, made it by reducing nitric oxide with iron filings in the presence of water. The most convenient method of prepara- tion is to heat ammonium nitrate cautiously, and to collect the gas over water: NH 4 NO 3 = N 2 O + 2H 2 O Nitrous oxide is a colorless gas; its solubility at is 130 vol- umes in 100 volumes of water, and at 25, 60 in 100; when liquefied it boils at -89.8. The compressed gas is furnished in cylinders and is used as an anesthetic. The discovery that nitrous oxide produces unconsciousness, or when breathed in smaller quantities causes hysterical laughing, was made in 1774 by Humphrey Davy. " Laughing gas " has been for a long time the " gas " adminis- tered by dentists. A mixture of nitrous oxide and oxygen is generally used, since by varying the proportions of the two gases a mixture can be attained which produces the desired degree of insensibility. Nitrous oxide is an endothermic compound; when it decom- poses 18,000 calories are liberated by each gram-molecule of the substance. The decomposition can be brought about by heat alone or better by a detonator, such as mercury fulminate. It will be recalled that ammonium nitrate is a powerful explosive when detonated; its use for this purpose is based upon the fact that when it decomposes, nitrogen and oxygen are set free with the evo- lution of a large amount of heat. Ammonium nitrate can be broken down into nitrous oxide and water by gentle heat without explosion; it is evident that the explosive property of the salt is due to the breaking apart of the nitrogen and oxygen in the nitrous oxide formed from it. 377. Nitrous oxide decomposes at the temperature at which most substances burn, and as oxj^gen is liberated the gas supports combustion. When nitrous oxide decomposes into its elements 2 volumes of the gas yield 2 volumes of nitrogen and 1 volume of oxygen: . 2N 2 O = 2N 2 + O 2 Since the product is one-third oxygen, whereas air contains but NITRIC ACID, NITROUS ACID, OXIDES OF NITROGEN 339 one-fifth oxygen, it is clear why substances burn, in general, with a more brilliant flame in nitrous oxide than in air. The proportion of oxygen is higher and the amount of inert gas to be heated is consequently less; and the decomposition of the nitrous oxide produces a large amount of heat. These factors lead to the pro- duction of a higher temperature, which results in increased lumi- nosity of the glowing gases in the flame. Nitrous oxide does not lose its oxygen at ordinary temperatures and, therefore, it does not affect metals as oxygen does; iron will not rust in the moist oxide as it does in air. For the same reason nitrous oxide does not convert nitric oxide into nitrogen dioxide at ordinary temperatures. It will be recalled that oxygen will bring about this reaction. We could readily distinguish nitrous oxide from air or oxygen by mixing it with some nitric oxide; if free oxygen is present the colorless nitric oxide is converted into the dioxide, which has a yellow-brown color. 378. We see from the behavior of nitrous oxide that an endo- thermic compound may be stable although it contains a large amount of bound up chemical energy. In order to make this energy available it is necessary to bring about the conditions which result in the disruption of the molecules. We can do this by raising the temperature of the compound. We know that we heat substances to make them interact, and it is reasonable to believe that before the atoms rearrange themselves to form new combinations it is first necessary to overcome the forces holding the atoms together in the molecule. As the temperature is raised the motion of the atoms in the molecules increases until finally a condition is reached which leads to the separation of the atoms; and when this occurs heat is given off or absorbed, as the case may be. The temperature at which decomposition first takes place varies with different substances. The study of detonators has brought out the fact that vibrations set up by one substance undergoing decomposition may induce similar vibrations within the molecules of another substance and thus cause the latter to decompose into its elements. This cannot take place, of course, when the decomposition of the molecule into its elements is associated with the absorption of energy. 379. Nitric Oxide. The importance of nitric oxide in chemical industry has been repeatedly emphasized; it serves as the cata- 340 INORGANIC CHEMISTRY FOR COLLEGES lytic agent in the chamber process for sulphuric acid, and its prepa- ration is the first step in the manufacture of nitric acid from nitro- gen. The gas is formed when nitric acid is brought into contact with substances which it can oxidize, and has been known for a long time. Priestley tested air by mixing nitric oxide with it; it united with the oxygen present and formed nitrogen dioxide, which could be dissolved in a solution of sodium hydroxide. It was, therefore, possible by measuring the decrease in volume of the air to find out how much oxygen it contained. Nitric oxide is most readily prepared by treating copper with concentrated nitric acid which has been diluted with an equal vol- ume of water. If the acid is not diluted, a part of the nitric oxide formed is oxidized by it to nitrogen dioxide. The equation for the reaction has already been discussed in detail (363) ; it is as follows : 3Cu + 8HNO 3 = 3Cu(NO 3 ) 2 + 2NO + 4H 2 9 The oxide prepared in this way is not pure, because it contains more or less nitrous oxide. It can be separated from the latter by passing it into a saturated solution of ferrous sulphate. The nitric oxide forms a compound with the sulphate and is, there- fore, absorbed by the solution. The brown nitrosyl ferrous sul- phate is the substance formed in the test for nitric acid (368). When the solution is heated nitric oxide is evolved. 380. Nitric oxide is a colorless gas, which is liquefied with dif- ficulty and is very slightly soluble in water. The liquefied gas boils at 153.6. The thermochemistry of nitric oxide has been discussed in some detail (358). It is an endothermic compound which contains a large amount of bound-up chemical energy. A comparison of the thermochemical equations for the decompo- sition of nitrous oxide and nitric oxide is of interest: 2N 2 O = 2N 2 + O 2 + 36,000 cals. 2NO = N 2 + O 2 + 43,200 cals. Like nitrous oxide, nitric oxide can be detonated and an explosion results. It is more stable than nitrous oxide. A higher tempera- ture is required to decompose it, and, as a consequence, it does not support combustion so readily. If a burning candle is introduced into the gas it is extinguished; nor will sulphur burn in the oxide. NITRIC ACID, NITROUS ACID, OXIDES OF NITROGEN 341 If, however, a bit of phosphorus which is actively burning is used, the combustion continues. In this case the temperature of the flame, which is higher, is sufficient to bring about the decomposi- tion of the nitric oxide. Nitric oxide takes on oxygen with great readiness, as we have seen. When the two gases are brought into contact at ordinary temperatures nitrogen dioxide, a brown gas, is formed : 2NO + O 2 <= 2NO 2 Nitric oxide unites directly with other elements than oxygen and with salts. It forms nitrosyl chloride (363) with chlorine, 2NO + C1 2 = 2NOC1, and nitrosyl ferrous sulphate, FeSO 4 ,NO, with ferrous sulphate. All these compounds decompose more or less readily into their constituents. 381. Nitrogen Dioxide and Nitrogen Tetroxide. The color of nitrogen dioxide changes markedly with change in temperature. At 150 it is deep brown, and as the temperature falls the color decreases in intensity. The gas condenses to a bright yellow liquid at 26; the color of this slowly fades with falling temperature and at 9 a white crystalline solid is formed. The study of this phenomenon has led to the conclusion that what we have called nitrogen dioxide is at ordinary temperatures a mixture of two compounds, one brown and one colorless. With rise in temperature the colorless compound is slowly converted into the brown oxide and with falling temperature the reverse phenomenon takes place; there is an equilibrium, markedly affected by temperature, which is represented by the following equation: N 2 O 4 <=* 2NO 2 - 13,600 cal. The brown gas is nitrogen dioxide and the colorless one is the polymer nitrogen tetroxide. This explanation of the change in color has been arrived at by determining the average molecular weight of the gas at different temperatures. Nitrogen dioxide can be made by the action of oxygen or air on nitric oxide or by heating nitrates of the heavy metals. Copper nitrate, as we have seen, yields the gas when heated (367) : 2Cu(NO 3 ) 2 = 2CuO + 4N0 2 + 2 342 INORGANIC CHEMISTRY FOR COLLEGES The gas is collected by passing it through a vessel kept cold by a mixture of ice and salt; it condenses to a yellow liquid. Nitrogen dioxide is an active oxidizing agent. Fuming nitric acid which contains dissolved nitrogen dioxide is a more powerful oxidizing agent than nitric acid itself. 382. Nitrogen dioxide can be detonated, and since 1 mole- cule furnishes 2 oxygen atoms which can be utilized for oxidizing other elements, a mixture of the liquid dioxide and compounds containing carbon and hydrogen is a powerful explosive. Per- fectly dry nitrogen dioxide does not affect benzene or naphthalene, but when the mixture is detonated an exceedingly violent explosion results. The force of the explosion is traceable to the heat lib- erated as the result of the separation of the oxide into its elements and that produced as a consequence of the union of the oxygen with the carbon and hydrogen present in the benzene or other sub- stance used. Only gases are formed, and as these are produced so suddenly that the heat generated cannot escape, a tremendous pressure is produced. Explosives based on this principle were used in the recent war in depth bombs. Nitrogen tetroxide reacts with cold water as follows : N 2 O 4 + H 2 O = HNO 2 + HNO 3 If the gas is passed into hot water the nitrous acid formed decom- poses as indicated by the following equation: 3HNO 2 = HNO 3 + 2NO + H 2 O By combining this equation with the one given above we have the following : . 3N 2 O 4 + 2H 2 O = 4HNO 3 + 2NO 383. Nitrous Anhydride. When a nitrite is treated in the cold with an acid, the nitrous acid first formed spontaneously decom- poses and forms a brown gas, which can be condensed at 21 to a blue liquid that has the composition represented by the for- mula N 2 O 3 : 2HN0 2 = H 2 O + N 2 O 3 The gas reacts with solutions of bases in the cold and forms nitrites. For these reasons nitrogen trioxide is usually called nitrous anhy- NITRIC ACID, NITROUS ACID, OXIDES OF NITROGEN 343 dride. The compound is very unstable and at its boiling-point begins to decompose into nitric oxide and nitrogen dioxide: N 2 3 <= NO + N0 2 At ordinary temperatures the dissociation is almost complete, and, consequently, when we pass the gas into a solution of a base we really have the simultaneous action of nitric oxide and nitrogen dioxide on it. 384. Nitric Anhydride. Nitric acid does not break down spon- taneously into its anhydride and water, but if it is treated with a powerful dehydrating agent the reaction takes place. Phosphorus pentoxide, which can be used for this purpose, unites with the water withdrawn and is converted into metaphosphoric acid: 2HN0 3 + P 2 5 = N 2 5 + 2HP0 3 When the mixture is cautiously heated, a brown gas is formed, which condenses to a liquid that boils at 45 and freezes to a white solid at 30. The anhydride unites with water and forms nitric acid, and decomposes slowly into nitrogen dioxide and oxygen: 2N 2 O 5 = 4N0 2 + O 2 EXERCISES 1 . What weight of Chile saltpeter containing 90 per cent NaNO 2 is required to prepare 1 ton of concentrated nitric acid, sp. gr. 1.41, provided the process is carried out in such a way that 95 per cent of the nitric acid theo- retically obtainable from the nitrate is obtained? 2. Write equations for the oxidation by nitric acid of the following: (a) Na 2 SO 3 , (6) As 2 O 3 to As 2 O 6 , (c) P to P 2 O 6; (d) CuHaOu to CO 2 and H 2 O, (e) Fe to Fe(NO 3 ) 3 . 3. (a) Calculate from the equation representing the reaction between hydrochloric acid and nitric acid the relation between the weights of the acids required in making aqua regia. (6) If concentrated hydrochloric acid (sp. gr. 1.2, 40 per cent HC1) and concentrated nitric acid (sp. gr. 1.4, 68 per cent HNO 3 ) are used, what is the relation between the weights of these acids required? (c) What is the relation between the volumes of the concentrated acids required? 4. Can HNO 3 be used to prepare (a) HC1 from NaCl, and (6) CO, from Na 2 C0 3 ? 5. When HNO 3 is made from NO prepared from N 2 and O 2 a large amount of energy must be furnished in preparing the oxide. When the acid is pre- pared by oxidizing NH 3 , the NO first formed is the result of an exothermic 344 INORGANIC CHEMISTRY FOR COLLEGES reaction; the union of N 2 and Ha also produces heat. At what stage in the second synthesis is the energy furnished? 6. Why is concentrated H 2 SO 4 mixed with HNO 3 in making cellulose nitrate? 7. Could carbon be used in preparing KNO 2 from KNO 3 ? Give a reason for your answer. 8. Explain why HNOa is a more active oxidizing agent than H 2 SO4. 9. How would you distinguish from one another the following: (a) NH 4 NO 3 , (6) KNO 2 , and (c) KNO 3 ? 10. How could you determine whether a sample of N 2 O contained a trace of NO? 11. How could a solution of HNO 3 free from HNO> be made from NO? Should the water used be hot or cold? Why? 12. In the reversible reaction N 2 O 4 <=^ 2NO 2 in which direction is the reaction exothermic? How would the equilibrium be affected by rise in temperature? 13. How would increased pressure affect the dissociation of nitrogen tetroxide when it is heated? 14. Would you expect that N 2 O 3 obeys (a) Boyle's law and (6) Charles' law? Give a reason for your answer in each case. 15. Chile saltpeter from which nitric acid is made contains a small percent- age of sodium chloride. Name two impurities that are introduced into the acid as a result of the presence of the salt. How could these be removed from the nitric acid? 16. When nitric acid is prepared it has a brown color due to the presence of NOa, How could this color be removed? CHAPTER XXIV THE DETERMINATION OF ATOMIC AND MOLECULAR WEIGHTS 385. Up to this point the student has gained a sufficient familiarity with chemical facts to make a consideration of the methods of determining atomic weights profitable. We have already seen that the law of definite proportions and the law of multiple proportions led Dalton to study the problem of deter- mining the relative weights of atoms, but as he was working with compounds, and, therefore, molecules, and had no way of deter- mining the number of atoms in the molecules it was impossible for him to arrive at figures that expressed the relative weights of the atoms themselves. Dalton found, for example, that water is made up by weight of 1 part of hydrogen and 8 parts of oyxgen. If a molecule of water contains 1 atom of each element and its formula is HO, then the atomic weight of oxygen is 8 if that of hydrogen is 1. If, however, the molecule contains 2 atoms of hydrogen and 1 of oxygen, H2O, the atomic weight of oxygen is 16; and other atomic proportions lead to different atomic weights for oxygen. 386. The Law of Gay-Lussac. In 1808 Gay-Lussac published his law of combining volumes, which states that when substances interact in the gaseous condition the volumes of the reacting substances and those of the products formed are in the relation of small whole numbers. For example, 2 volumes of hydrogen unite with 1 volume of oxygen and form 2 volumes of steam; 1 of hydro- gen and 1 of chlorine form 2 of hydrogen chloride, etc. The striking siniDlicity in the weight relations in chemical reactions summarized in the law of multiple proportions, has its counterpart in the equally striking simplicity in the volume rela- tions observed when reaction takes place between substances in the gaseous state. We have seen how an attempt to give a physical 345 346 INORGANIC CHEMISTRY FOR COLLEGES interpretation of the law in regard to weight relations led to an atomic conception of matter which was susceptible of experi- mental investigation and, therefore, unlike the old Greek views; and we shall now see how the law summarizing volume relations, as the result of a similar endeavor, furnished the key to the prob- lem of determining the relative weights of molecules and atoms. The law of multiple proportions has to do with the relationship between atoms; the law of Gay-Lussac summarizes the relation- ship between molecules. And we have just seen that Dalton's attempt to determine atomic weights was futile because he had no way of determining the number of atoms in molecules. The significance of Gay-Lussac's law was first seen by Avo- gadro, an Italian, in 1811, but the important conclusions which he drew from it were not appreciated until later, when Cannizzaro, in 1860, demonstrated how Avogadro's conception furnished a satisfactory method upon which to base the determination of atomic weights. For years chemists had been changing from one system to another, and a number of these were in use; as a conse- quence, one chemist wrote the formula H^O for water and another HO. The system of atomic weights used to-day is based on the method of determination outlined by Cannizzaro, who was able to clear up the difficulties existing because he possessed the power to marshal the facts in such a simple and logical way that their sig- nificance was understood. 387. The Law of Avogadro. Avogadro clearly differentiated atoms from molecules the molecule was the smallest physical unit of a substance and was made up of atoms. And on account of the simplicity of the volume relations involved when molecules interact he emphasized the necessity of considering these relations because they were without doubt related to the volumes of the molecules themselves. Let us see how a study of the volumes occupied by gaseous substances helps in the problem which baffled Dalton. In the selection of a standard of atomic weights the logical thing to do is to select the lightest element known and assign to it the atomic weight 1; this was done and hydrogen was selected (66). It would be interesting at the outset to compare the volumes occu- pied by a number of gaseous substances, and to have a basis for the comparison we shall take the amount of each gas which con- DETERMINATION OF ATOMIC AND MOLECULAR WEIGHTS 347 tains 1 gram of hydrogen. These volumes are recorded in the following table: Element combined with hydrogen Cl o N c S P Si H Volume in liters at and 760 mm. containing 1 gram of hydrogen Relative volumes with largest as unity 22.4 1. 11.2 i 7.5 A 5.5 | 11.2 i 7.5 \ 5.5 i 11.2 A In the last column is given hydrogen gas. The simple relation that exists between the volumes of the gases listed which contain 1 gram of hydrogen is striking; and we would obtain similar results if other hydrogen compounds were studied. A natural question is What is the physical basis for this? It is hardly possible that the molecules of the different gases all contain the same number of hydrogen atoms. What would be the result if they differed in this respect? If we prepare from the same amount of hydrogen in each case molecules which contain respectively 1, 2, 3, and 4 atoms of hydrogen, the number of molecules formed in the second case will be one-half as many as those formed in the first case; the number formed in the third case will be one-third as many as in the first case; and the number formed in the fourth case one-fourth as many. This fact when considered along with the fact that the volumes of the gases formed from the fixed weight of hydrogen are in the relation of 1 to i, J, and J leads to the reasonable assumption that the volume of a gas is determined by the number of molecules it contains, whatever the molecule may be. This conclusion is known as Avogadro's Law and is usually stated as follows : Equal volumes of all gases at the same temperature and pressure contain the same number of molecules. For years this remarkable conclusion of Avogadro was known as a hypothesis for no one was able to count the number of mole- cules of a gas; and it seemed impossible that physical experimen- tation could ever reach such a point of refinement that it could deal with particles so small that billions of billions of them could be contained in a space the size of a drop of water. The hypothesis became, however, the basis of the determination of atomic weights 348 INORGANIC CHEMISTRY FOR COLLEGES and, therefore, had an important significance; it was, accordingly, known later as Avogadro's rule. Recent work in physics has made it possible to count molecules, and as a number of independent methods all lead to the same conclusion and confirm the hypoth- esis, it is now known as a law. One gram of hydrogen contains 3 X 10 23 molecules. 1 388. The Determination of Atomic Weights. Let us now see how the fact that equal volumes of gases contain the same number of molecules can be used in building up a system of atomic weights. We will take as our standard hydrogen the lightest known sub- stance and call its atomic weight 1. The volumes of as many compounds of hydrogen as possible are determined, selecting in each case the volume which contains 1 gram of hydrogen. The results in a number of cases are given in the table already dis- cussed (page 347). It is evident that we shall get the largest volume in the case of the gas that has but 1 atom in the molecule for in this gas we shall have the largest number of molecules. Accordingly, in the compound which occupies the largest volume we assume that 1 atom of hydrogen is present in the molecule. This assumption is justified when we select the compound contain- ing chlorine, hydrochloric acid, for there is no gas known which has a larger volume than 22.4 liters when the amount of gas selected is that which contains 1 gram of hydrogen. These considerations lead us to the view that the molecule of hydrochloric acid contains 1 hydrogen atom. We can determine the number of chlorine atoms in the molecule in the same way by examining the volumes of chlorine compounds all of which contain the same weight of the element. We find in this case that none of these yields a greater volume than the compound containing hydrogen, and we con- clude, therefore, that hydrochloric acid contains 1 chlorine atom. If there is 1 atom of hydrogen and 1 atom of chlorine in hydro- chloric acid its formula is HC1. We find by experiment that 1 gram of hydrogen unites with 35.5 grams of chlorine and forms 36.5 grams of hydrochloric acid; the atomic weight of chlorine is, therefore, 35.5, and the molecular weight of hydrochloric acid is 36.5. We also find by experiment that 36.5 grams of hydrochloric acid 1 gram-molecular-weight or mol occupies 22.4 liters. 1 This is a short way of expressing the number made up of 3 followed by 23 zeros; it is 300,000 billion billions. DETERMINATION OF ATOMIC AND MOLECULAR WEIGHTS 349 We have thus a definite experimental basis founded on fact for the determination of the atomic weight of chlorine and the formula of hydrochloric acid; and the same method can be applied in other cases. The only thing that would lead to a change in the atomic weight of chlorine would be the discovery of a compound which occupied as a gas more than 22.4 liters, when the weight of it con- taining 1 gram of hydrogen was used; for the gas which occupies the largest volume has 1 atom, and in this case there would be more than 1 hydrogen atom in hydrochloric acid. The method used in determining atomic weights can be further emphasized by applying it to oxygen. The same reasoning is used. The volumes of gaseous compounds containing oxygen are deter- mined, using in each case the amounts of the several substances which contain the same weight of oxygen. Results which have been obtained in a number of compounds are given in the following table: Element combined with oxygen H as C as Gas Sas Sas Nas Nas Vol. of substance at and 760 mm. contain- ing 1 gram oxygen H 2 O 1.4 CO 1.4 CO 2 0.7 SO 2 0.7 SO 3 0.48 NO 1.4 NO 2 0.7 Relative volumes with largest as unity 1 1 | | i 1 i From these results we draw the conclusion that water, carbon monoxide, and nitric oxide each contains 1 atom of oxygen. An examination of the table on page 347 will give us information as to the number of hydrogen atoms in water. The volume of the water-vapor which contains 1 gram of hydrogen is one-half the volume of the hydrochloric acid containing this weight of hydrogen, and, therefore, there are twice as many atoms of hydrogen in water as in hydrochloric acid, and its formula is H^O. Since 18 grams of water contain 2 grams of hydrogen and 16 grams of oxygen, the atomic weight of oxygen is 16 when hydrogen is 1. 389. The Determination of Formulas. Having arrived in this way at the atomic weights of the elements we can readily deter- mine the formula of any compound. Avogadro's law tells us that 350 INORGANIC CHEMISTRY FOR COLLEGES equal volumes of all gases contain the same number of molecules; as a consequence, the weights of equal volumes of any two gases are in the same relation as the weights of the individual molecules of these gases. We have seen that 22.4 liters of hydrochloric acid weigh 36.5 grams and that the molecular weight of hydro- chloric acid is 36.5. If, therefore, we determine the weight in grams of 22.4 liters of any gas the number obtained will be the molecular weight of the gas. The volume 22.4 liters is, there- fore, one of the greatest significance in chemistry; for it is the volume of a gram-molecular-weight of all gases. We can deter- mine, therefore, the molecular weight of any gas by weighing a sample of known volume and calculating from the result the weight of 22.4 liters the number obtained is the molecular weight of the gas. In order to write the formula of a substance we must know, in addition to the weight of its molecule, the proportion by weight of the elements of which it is composed, and the atomic weights of these elements. An example will show how this can be done. A substance was found upon analysis to contain 30.43 per cent nitrogen and 69.57 per cent oxygen; the weight of 22.4 liters at and 760 mm. was 46 grams. What is its formula? The molecular weight is 46; 30.43 per cent of this is nitrogen, there- fore, 46 X .3043 = 14 is the weight of the nitrogen in the mole- cule. The weight of the oxygen is 46 X .6957 = 32. Since an atom of nitrogen weighs 14 and 1 of oxygen weighs 16 the for- mula is, evidently, NO2. The formula of any compound can be determined in this way. The weight of 22.4 liters at and 760 mm. of the substance in the form of a gas is multiplied in turn by the percentage of each element present, and the result in each case divided by the atomic weight of the element. The numbers obtained are the numbers of the several atoms present in the molecule. 390. The formulas of elementary gases can be determined in the same way. For example, 22.4 liters of hydrogen weigh 2 grams. Since the atomic weight of hydrogen is 1 the formula of hydrogen gas is Eb. It is, of course, not necessary for the substance to exist as a gas at in order to get its molecular weight in the gaseous con- dition. The volume of a given weight is determined at any con- DETERMINATION OF ATOMIC AND MOLECULAR WEIGHTS 351 venient temperature and pressure and a calculation is made by applying the gas laws to determine what the volume would be at and 760 mm. In this way the molecular weight of phosphorus in the gaseous condition has been determined although the element is a solid at ordinary temperatures. Since the molecular weight was found to be 124 and the atomic weight is 31 the molecule in the gaseous condition at the temperature used is P4. 391. Atomic Weights Based on Oxygen as Standard. The system of atomic weights which has been explained is based on the hydrogen atom to which was assigned the value 1, and round numbers have been used in order to simplify the discussion. The most careful analyses of water that have been made lead to the conclusion that 2 grams of hydrogen unite with 15.88 grams of oxygen and, as a consequence, the atomic weight of oxygen is 15.88 if that of hydrogen is 1. Using this value the atomic weight of other elements can be determined; for example, the weight of copper that unites with 15.88 grams of oxygen to form copper oxide, CuO, is the atomic weight of copper, namely 63.09. The atomic weights of a large number of elements were found by determining the relative weights of the element and oxygen in their oxides; and the values of the atomic weights were calculated with the aid of the atomic weight of oxygen. The value of the latter, as we have seen, is deduced from the results of the analysis of water, and the actual numbers used for the atomic weights of all elements calculated from the analysis of their oxides is deter- mined, therefore, by the figures obtained in the study of the com- position of water. As methods of analysis increased in accuracy as balances were improved, for example the results of the analysis of water gave different and more accurate figures, and the atomic weight assigned to oxygen kept changing. For a long time 15.79 was accepted as the atomic weight of oxygen and the atomic weight of other elements calculated from this figure. Then more accurate work showed that 2 grams of hydrogen united not with 15.79 grams of oxygen, but with 15.88 grams; and 15.88 was taken as the atomic weight of the element. As a result, the atomic weights of all elements which were determined by analyzing oxides had to be recalculated and new figures were obtained. Every time the atomic weight of oxygen was changed, as a result of a more accurate determination of the ratio between the weights of 352 INORGANIC CHEMISTRY FOR COLLEGES hydrogen and oxygen that unite, all atomic weights had to be recal- culated. This was an unfortunate state of affairs and resulted largely from the fact that the accurate determination of the ratio between the weights of hydrogen and oxygen that unite to form water is one of the most difficult of analytical processes, because hydrogen is the lightest substance known and the weighing and measuring of large volumes of a gas are exceedingly difficult operations. To see what could be done to avoid this constant changing of atomic weights an international commission of chem- ists was appointed to devise a plan to be followed. This resulted in taking as the standard oxygen and calling its atomic weight 16. As a result, the atomic weight of hydrogen became 1.008, and any changes in the oxygen-hydrogen ratio affected the atomic weight of hydrogen alone. The value accepted for the atomic weight of copper, for example, will be changed only when more accurate analyses of copper compounds lead to different results from those obtained in the past. 392. The Determination of Atomic Weights by Analysis. It has been pointed out that the accurate measurement of gases is difficult, and, as a consequence, the values of the atomic weights obtained in this way are not the most accurate that can be arrived at. We determine, for example, in the way described, from gases containing carbon, that its atomic weight is approximately 12, but we arrive at a more accurate figure by using solid or liquid compounds containing the element. A method that could be used will illustrate the procedure. A carefully weighed amount of carbon could be burned in a stream of oxygen and the gas formed, CO2, absorbed by sodium hydroxide. The increase in weight of the vessel containing the sodium hydroxide gives the weight of carbon dioxide produced. The difference between this weight and the weight of carbon burned is the weight of the oxygen. From the weight of carbon and the weight of oxygen and knowing the formula of carbon dioxide and the atomic weight of oxygen, the atomic weight of carbon can be calculated. 393. The Law of Dulong and Petit. The system of determining atomic weights which has been described is based on conclusions drawn from a study of the volumes of gaseous compounds of the elements. If such compounds do not exist, we must employ some other method which yields results that fit into the system in use. DETERMINATION OF ATOMIC AND MOLECULAR WEIGHTS 353 In general, the acid-forming elements yield compounds that are gases or can be readily vaporized, while most metals do not form such compounds. If a method could be found to determine the atomic weights of metals which did not involve the use of gases it could be linked up with the method that has been described, because some metals, such as mercury, do form compounds which can be vaporized. A determination of the atomic weight of such a metal by the two methods should give the same result if the one applied to metallic atoms is to be used; for it is necessary to have the same standard for all atomic weights. All gases have the same molecular volume that is, a molecule of any gas occupies the same space as a molecule of any other gas; and we have used this fact to build up a system of atomic weights. There is no such generalization in the case of liquids and solids, and some other property than the relation between weight and volume must be used in deciding upon the atomic weights of metals. Dulong and Petit, two French chemists, investigated the specific heats of solid elements and discovered an important generalization which was announced in 1819. It will be recalled that the specific heat of a substance is the number of calories required to raise the temperature of 1 gram of it 1 degree. The specific heats of elements vary widely; a few values for metals are as follows: iron, 0.112; zinc, 0.093; gold, 0.032. It was noticed that in a series of metals as the atomic weight increased the specific heat decreased; and the investigators were led to multiply these values in the case of each element. The products obtained were approximately the same number. For example, iron, atomic weight 56, specific heat 0.112, 56 X 0.112 = 6.3; gold 197 X 0.032 = 6.3; zinc 65.4 X 0.093 = 6.1. The average value of the product in the case of metals is 6.4. Such facts as these led to what is called Dulong and Petit's law, which states that the product obtained by multiplying the specific heat of a metal by its atomic weight is a constant in all cases. The physical significance of this is clear. The product obtained in each case, approximately 6.4 calories, is the heat required to raise 1 gram-atomic-weight of the metal 1 degree. So the law can be stated in the following form: " The atoms of all metals have the same heat capacity." This generalization will, accordingly, give us a means of determining the relative atomic weights ; all that is needed is a standard, and this must be the same as that used in 354 INORGANIC CHEMISTRY FOR COLLEGES the case of the gaseous elements. It is well to point out here again that the system of determining the atomic weights of elements that exist in gaseous compounds is based on the fact that all molecules occupy the same volume; and we shall see that the system for obtaining the atomic weights of solid elements is based upon the fact that the atoms of all metals have the same capacity for heat. In order to determine the atomic weight of a metal all that is necessary is to determine its specific heat and make use of the fact that atomic weight X specific heat = 6.4 If the specific heat of a newly discovered metal were determined to be 0.04 its atomic weight would be found as follows: at. wt. X 0.04 = 6.4 at. wt. = 160 The application of the law of Dulong and Petit yields approxi- mate values only for the atomic weights because the product of the specific heat and the atomic weight is not exactly the same number in all cases. The reasons for these variations were un- known for a long time. Recently, however, the advance in our knowledge of energy and its transfer has helped in interpreting the deviations from the law of Dulong and Petit. Extension of our knowledge of matter has frequently resulted from a search for a reason for the fact that the so-called constant in a law varies slightly when we apply the law to specific cases. The deviations are often very small, and they are discovered only when measurements are made with the greatest care a fact which emphasizes again the value of the most accurate work possible in studying the quantity relationships in physical and chemical phenomena. By the use of the law of Dulong and Petit the approximate value of the atomic weight of a metal can be determined ; the accurate value is obtained by the analysis of a compound containing the element. For example, we calculate from the specific heat of copper that its atomic weight is approximately 63. By a careful analysis of copper oxide we find that 63.57 grams of copper are combined with 16 grams of oxygen; the atomic weight of the element is, therefore, 63.57. DETERMINATION OF ATOMIC AND MOLECULAR WEIGHTS 355 The law of Dulong and Petit gives us a way to determine the relative atomic weights of the metals, but it does not fix a standard. In order to get the constant to be used in calculating the atomic weights of elements from their specific heat, we must multiply the specific heat of some element by its atomic weight we must have a standard. Fortunately some metals form volatile compounds and we can determine the atomic weights of such metals with the aid of ttn method applied to gases and obtain values based on the standard oxygen 16 used in the case of the non-metallic ele- ments. The law of Dulong and Petit, we have seen, leads to the con- clusion that the same amount of heat is required to cause the same rise in temperature of the atoms of all metals. This is a most remarkable fact; we would not expect that the same amount of heat would produce approximately the same change in temperature in the case of a light atom like sodium, which weighs 23, and a heavy atom like lead, which weighs 207. As a result of the appli- cation of recent theories in physics to the study of atoms, new light is being thrown on this important fact. In the whole field of chemistry there is opportunity for research; we are just at the threshold of the science. 394. The Importance of Accurate Atomic Weights. The accurate determination of the atomic weights of the elements is one of the most important problems in chemistry, and great ingenuity has been shown in increasing the accuracy of methods of analysis used in their determinations. The atom is the chemical unit of matter, and the more certain our knowledge is of its prop- erties the more fully we understand the changes in matter and the better we can make use of them. The weight of an atom is but one of its properties, but it is an important one; one of the most fundamental generalizations of chemistry the periodic law is based on the atomic weights, and this could have been discovered only after reasonably correct values of the atomic weights were known. It is highly probable that as the knowledge of the true atomic weights becomes more exact other generalizations will be apparent, and causes for phenomena not yet understood will become evident. The determinations of the atomic weights of some of the ele- ments are among the most accurate measurements which have ever 356 INORGANIC CHEMISTRY FOR COLLEGES been made in science. A number of chemists specially equipped for this kind of work, which requires physical dexterity, infinite patience, the ability to see hidden sources of error, and the ingenu- ity necessary to invent new and more accurate methods of measure- ment, have obtained the admirable results summarized in a table of atomic weights. Stas, a Belgian chemist (1813-1891), set an example to the chemical world by his work in determining atomic weights. His results were accepted as the standard of accuracy until revised later by Professor T. W. Richards of Harvard Uni- versity, who, with his co-workers, has obtained values for the atomic weights of the more important elements, which surpass in accuracy all previous determinations. Professor Edward Morley, formerly of Adelbert College, spent a number of years studying the hydrogen-oxygen ratio, and his results are universally accepted as the most accurate that have been obtained for this constant. It is necessary that chemists the world over agree to use the same values for the atomic weights, and as these numbers are con- stantly being revised, an international committee formed of rep- resentatives from the chemical societies of the world studies the new work in this field and publishes each year a table of atomic weights. Before there was an international agreement as to the value of atomic weights difficulties arose which in some cases affected commercial relations. At one time, for example, the accepted atomic weight for chromium was not the same in Eng- land and America. In calculating the amount of chromium in an ore from its analysis, it is necessary to use the atomic weight of the element, and the amount found, and consequently the value of the ore will be different if different atomic weights are used. It is necessary, therefore, in commercial transactions of this kind to agree on atomic weights, and the values selected by the inter- national commission are accepted and used in all calculations. EXERCISES 1. A number of compounds gave the following results. Calculate the formula in each case: (a) 91.30 per cent P and 8.65 per cent H 2 ; 1 liter weighed 1.52 g. (6) 87.59 per cent Si and 12.45 per cent H 2 ; 1 liter weighed 1.44 grams, (c) 80.20 per cent C and 20.1 per cent H 2 ; 1 liter weighed 1.33 grams, (d) 92.36 per cent C and 7.8 per cent H 2 ; 1 liter weighed 1.16 grams. 2. An experiment showed that when 1.2714- grams of copper were heated DETERMINA TION OF A TOM 1C AND MOLECULAR WEIGHTS 3 57 in a stream of air 1.5920 grams of CuO were formed. Calculate the atomic weight of Cu, assuming the formula of the oxide to be CuO. 3. In a determination of the atomic weight of bromine 2.6970 grams of silver were dissolved in HNO 3 and precipitated as AgBr; the weight of the latter obtained was 4.6950 grams. Calculate the atomic weight of bromine, using 107.88 as the atomic weight of silver. 4. In determining the atomic weight of sodium 1.2356 grams of pure NaCl were dissolved in water and treated with a solution of silver nitrate. The silver chloride obtained weighed 3.0274 grams. Taking 107.88 as the atomic weight of silver and 35.46 as that of chlorine calculate the atomic weight of sodium. 5. The specific heat of a metal was found to be 0.033. When 5.0150 grams of the metal were converted into an oxide the latter weighed 5.4150 grams. Calculate the atomic weight of the element. 6. The specific heat of a metal was found to be 0.031. When 4.1440 grams of the metal were converted into its chloride the latter weighed 5.5624 grams. Calculate the atomic weight of the metal, using 35.46 as the atomic weight of chlorine. CHAPTER XXV THE PERIODIC LAW 395. The elements of which the world is made up show a great variety of physical and chemical properties. Some are light gases and some are dense solids; some, like sulphur, are exceedingly poor conductors of heat and electricity,, and others, like silver, are excellent conductors. There is just as wide a range in chemical properties ; we have active metals that form strong bases, of which sodium is an example, and equally active non-metals, like chlorine, that yield strong acids. Other elements are characterized by being excessively inert; nitrogen and carbon show little or no activity at ordinary temperatures. Elements exist that resemble each other strikingly in chemical properties. Calcium, for exam- ple, forms a series of compounds the chemical behavior of which is markedly like that of the compounds of magnesium of analogous composition. The question before the chemist has always been what is the fundamental cause underlying these differences and similarities. After Dalton had proposed the atomic theory and the atomic weights of a number of elements had been determined, it was noticed that there was some relation between these numbers and the properties of the elements. The elements chlorine, bromine, and iodine resemble one another markedly in chemical properties; they form with other elements compounds of analogous composi- tion, which resemble one another in their chemical behavior. When the physical properties of these elements and their com- pounds are studied it is found that there is a progressive change as we pass from chlorine to bromine to iodine. For example, chlorine is a yellow gas, bromine a red liquid, and iodine a black solid. The solubilities of the compounds of the formulas NaCl, NaBr, and Nal, are, respectively, 35.86, 88.76, and 177.9 grams in 100 c.c. 358 THE PERIODIC LAW 359 of water at 18. On account of this close relationship these ele- ments were said to belong to a chemical family. The atomic weights of chlorine, bromine, and iodine are, respectively, 35.46, 79.92, and 126.92. The increase in these values is progressive, and it is a remarkable fact that the atomic weight of bromine is not far from the mean of the values of the other two, which is 81.19. These facts are of interest, but become of great significance when considered along with the fact that other so-called families of elements exist, and that in these there is the same numerical relationship between the atomic weights of the members of any family. It was clear that there was some connection between the prop- erties of elements and their atomic weights, and a number of attempts, more or less successful, were made to correlate these facts. It was some time after they were known, however, before the real relationship was discerned. Mendelejeff, a Russian chem- ist, in an endeavor to classify the properties of elements and their compounds, arranged the former in the order of their atomic weights, and examined them from this point of view. He found that starting with the element with the smallest atomic weight, hydrogen, the chemical properties of the succeeding elements changed markedly as the atomic weight increased. Lithium, which has the atomic weight 6.9, is an active base-forming ele- ment which has the valence 1 ; the next element, beryllium, atomic weight 9.1, has less active base-forming properties and has the valence 2; boron, atomic weight 11, forms weak acids and has the valence 3; next came carbon, atomic weight 12, which is also an acid-forming element with the valence 4; then nitrogen, 14, with the valence 5, oxygen, 16, and, finally, fluorine, 19. The next element which was sodium, 23, resembled lithium markedly in chemical properties and had the same valence as the latter. Magnesium, which followed, was like beryllium, aluminium like boron, silicon like carbon, phosphorus like nitrogen, sulphur like oxygen, and chlorine like fluorine. Potassium, the element which came next, was the first member of a series of seven ele- ments the properties of which varied with increasing atomic weight in the same way as the properties of the seven elements in the series from lithium to fluorine, inclusive, and those in the series of seven elements from sodium to chlorine, inclusive. These facts 360 INORGANIC CHEMISTRY FOR COLLEGES become clearer if the symbols of the elements, exclusive of hydrogen, and their atomic weights are set down as follows : Li, 6.9 Be, 9.1 B, 11. C, 12. N, 14. O, 16. F, 19. Na, 23. Mg, 24.3 Al, 27.1 Si, 28.3 P, 31. S, 32. Cl, 35.5 K, 39.1 Ca, 40.1 Sc, 44.1 Ti, 48.1 V, 51. Cr, 52. Mn, 54.9 The element which follows manganese, the last element in the above tabulation, is iron, 55.9; it does not resemble lithium, sodium, and potassium. Nickel, 58.7, and cobalt, 59, are very much like iron in chemical and physical properties. Consequently, these three elements were set off by themselves and became members of the eighth group of elements, the groups consisting of the members in the vertical columns in the above tabulation; they are numbered 1 to 7 from lithium to fluorine. Copper, 63.6, which followed cobalt, became the first member of a new series of elements the seventh member of which was bromine, and, consequently, fell into its place under chlorine in the tabulation. The arrangement following this method was continued in the same way with the rest of the elements, and led to a table which resembled closely that given facing the inside of the back cover of this book. Men- delejeff put forward his classification of the elements in 1869 and did not include what is called group O in the table, as these elements, the so-called noble gases, were unknown at that time. From what has been brought out in the above discussion it is evident that there is a striking relationship between the physical and chemical properties of the elements and their atomic weights, or to use a mathematical term, the properties of the elements are a function of their atomic weights. They are not, however, a direct function of the atomic weights, for as the latter increase there is a recurrence in properties over and over again; for this reason they are said to be a periodic function. Mendelejeff's study of the subject led him to this conclusion, and when he published his first table of atomic weights he proposed the periodic law, which states that the physical and chemical properties of the elements are a periodic function of their atomic weights. 396. A few additional words will be necessary to complete the description of the table of atomic weights. In the first horizontal THE PERIODIC LAW 361 line under the numbers of the groups are given the formulas E2O, EO, etc. These are general formulas to represent the composition of the oxides of the members of the group ; for example, in group 3 we have boron, scandium, etc., the oxides of which have the general formula E 2 O 3 , that is, B 2 O 3 , Sc 2 O 3 , etc. In the second horizontal line, we have beginning with group 4 the symbols EH4, EH 3 , etc. These are general formulas for the hydrides of the elements in the several groups, such as CH4, NH 3 , etc. In the third line appear the letters A and B in each group. These serve to designate the sub-groups or families into which the mem- bers of the groups are divided. For example, while there is a general likeness in certain respects between all the members of group 2, the most striking similarities appear when we consider calcium, strontium, and barium on the one hand, and magnesium, zinc, cadmium, and mercury on the other. To emphasize this fact the symbols of one set of elements are placed on one side of the column in the tabulation and those of the other on the other side. The dots, . . . . , placed in some of the squares indicate the fact that elements to fit into these places have not yet been discovered. When Mendelejeff proposed this classification of the elements, the elements now known as scandium, group 3 series 3, gallium, group 3 series 4, and germanium, group 4 series 4, were unknown. If the tabulation had been completed according to the method being used, when titanium, the next known element after calcium, was reached, the regularities observed in the first part of the table would have disappeared. Mendelejeff saw that if the place under boron, which is now occupied by scandium, was left vacant, titanium would fit into the place under carbon and the regularities would not be disturbed. He did this and prophesied that an element with an atomic weight of approximately 44 would be discovered, and that in chemical properties and in the com- position of the compounds prepared from it, it would resemble boron. This discovery was subsequently made. Mendelejeff left vacant the spaces now occupied by gallium and germanium for the same reasons, and his prophecies in these two cases were also fulfilled. These facts confirmed in a striking way the truth under- lying the periodic classification and had a marked influence in leading chemists to recognize its value. 397. When the system was proposed the atomic weight accepted 362 INORGANIC CHEMISTRY FOR COLLEGES for uranium, group 6 series 10, was 119.2 one-half the number assigned to it to-day. This atomic weight would place the ele- ment in group 5 series 6 under arsenic. As it was evident that its properties were not in accord with this position, Mendelejeff maintained that its atomic weight should be doubled ; he did this and uranium fell into the place under molybdenum and tungsten, W, where it belonged. Subsequent investigation confirmed the correctness of Mendelejeff 's view. It was evident that the atomic weights accepted for some of the other elements were incorrect; these were placed in the table in accordance with their properties, and in nearly every case redetermination of the values led to a confirmation of the correctness of the conclusions indicated by the periodic law. In the case of tellurium, group 6 series 6, however, this has not been the case. It belongs, evidently, to the sulphur-selenium family and it is placed in the position to indi- cate this fact, although its atomic weight being greater than iodine, puts it in group 8. The periodic law has been a guide of the greatest value in chemical investigation for over fifty years; it has suggested the possibility of the preparation of many compounds with valuable properties, and, after their discovery, such compounds have been studied from the standpoint of the law. Its greatest value, how- ever, is, perhaps, the service rendered in assisting in the classifica- tion of the many facts of inorganic chemistry. The student will find it helpful in remembering many of these facts. Since science is systematized or classified knowledge, the periodic law has been a most important factor in the development of the science of chemistry. As the elements and their compounds have been intensively studied in recent years, many facts have been discovered which are in striking accord with the periodic classification; but others in an equally striking manner cannot be interpreted through the use of this generalization. These exceptions do not lead the chemist to refuse to accept the law as an expression of a truth, but rather inspire him to a more detailed study of the facts in the hope of discovering the causes of the exceptions, in order to arrive at a more perfect expression of the truth in a modified law. 398. The periodic law considers the elements solely from the standpoint of the weight of the atoms of the elements. It is THE PERIODIC LAW 363 evident that it must be an imperfect expression of the truth in regard to the properties of the elements if their content of energy is a factor in determining their properties. The study of the atom from this point of view began with the discovery of radium. A great deal of evidence has been found for the hypothesis that the atoms of the elements are not single units but are made up of so-called electrons which are charges of negative electricity that surround a nucleus having a positive charge. Since the atom contains positive and negative electricity it is a storehouse of energy ; and when it unites with another atom a part of this energy is transformed in the chemical union. According to this view the elements have been built up by progressively increasing the number of positive charges and negative charges which go to make up the atoms of the several elements. The periodic classification of the elements which results from the conception of the electrical constitution of the atom is more in accord with the facts than that arrived at from the consideration of the weight relations; the exceptions to the law of Mendelejeff disappear, and many facts which cannot be interpreted by the law or could not be foretold by it are in accordance with the newer classification (798) . The periodic classification of the elements cannot be advan- tageously discussed in detail without a knowledge of a large number of facts. For this reason it will be considered as the chemistry of the elements and their compounds is developed in subsequent chapters. Its application to the facts as they are brought forward will bring out clearly the importance of the law as an aid in their generalization. The classification according to Mendelejeff will be used, as it serves adequately in the consideration of the facts to be presented. CHAPTER XXVI THE HALOGEN FAMILY 399. The members of the so-called halogen family constitute sub-group B in the seventh group in the periodic classification of the elements. A study of the members of this sub-group from the point of view of searching out similarities and differences in chemical and physical properties, will bring out clearly the sig- nificance of the periodic law ; and the student will see more clearly than in the past how the many facts of chemistry may be corre- lated and thus become examples of general principles. A compari- son of the chemistry of bromine with the chemistry of chlorine will necessitate a review of the facts concerning the latter, which will thus be impressed on the memory more definitely. Repe- tition is the most important factor in memory, and when the repe- tition of a set of facts is associated with the acquisition of new facts of importance and with the discovery of relationships of interest, the process becomes a source of mental satisfaction. The student should strive to study in this way, and when the method has become a habit of thought as the result of repetition, one of the attributes of a trained mind will have been attained, and one of the objects of education accomplished. The members of the halogen family are active acid-forming, electro-negative elements. Their activity in this respect decreases with increasing atomic weight. This is evident from a consider- ation of the change in energy which takes place when the elements react with hydrogen, a typical electro-positive element. The heats of formation (165) of the halogen hydrides, HF, HC1, HBr, HI are, respectively, 38.5, 22, 8.5, and 0.5 (at 400) large calories. The values are those obtained in the production of 1 gram-molec- ular-weight of the hydrides from hydrogen and the halogens in the gaseous condition. These values indicate clearly the order in which the elements stand when arranged according to their 364 THE HALOGEN FAMILY 365 chemical activity. This order is just the reverse of that arrived at if the arrangement is made according to the weight relations. This means that the activity of the elements decreases as their atomic weights increase. This relationship is shown in the case of most families of elements the lighter atoms are the more active. We shall see as the subject develops, that if we arrange these elements or their compounds in which the former play the role of electro-negative elements, according to any property which results from chemical activity or from the weight of the atoms, the order will, in most cases, be the same, namely, fluorine, chlorine, bromine, and iodine. The facts about to be presented should be examined from this point of view, BROMINE 400. Occurrence and Discovery. Bromine is found chiefly as sodium bromide and magnesium bromide associated with the chlorides of these elements. As the salts are soluble they find their way into the ocean, and the chief sources of the halogen are salt mines that have, in all probability, been formed as the result of the drying up of inland seas. The quantity of bromides found is very small compared with that of the chlorides. The element was discovered by Balard in 1826 as the result of treating with chlorine a solution from which crude salt had been crystallized the so-called mother-liquor. The bromides present in the latter reacted with chlorine and formed chlorides, and bromine was set free. The reaction in the case of sodium bromide is as follows: 2NaBr + C1 2 = 2NaCl + Br 2 A deep yellow color was formed in the solution. Balard shook the latter with ether, which dissolved the colored material and thus separated it from the water. On evaporation of the ether the residue left was found to be a red liquid. On account of its marked odor the name selected for it was derived from the Greek word signifying a stench. Salt beds which contain bromine in the proportion which makes it profitable to use them as a source of bromine occur at Stassfurt in Germany, and in the States of Michigan, Ohio, and West Virginia. 366 INORGANIC CHEMISTRY FOR COLLEGES 401. Preparation of Bromine. In the manufacture of bromine on the industrial scale the reaction discovered by Balard is gen- erally used. The brine, obtained by dissolving the crude salt in water, is treated with just enough chlorine to liberate the bromine present. Steam is then passed through the liquid and the bro- mine passes off and is condensed along with the steam, and sep- arates as a heavy, reddish-brown liquid. Bromine can be separated from its compounds by utilizing the methods which have been described in connection with the prep- aration of chlorine. It is formed, for example, when manganese dioxide is treated with hydrobromic acid, MnO 2 + 4HBr = MnBr 2 + 2H 2 O + Br 2 and when a bromide is gently heated with manganese dioxide and concentrated sulphuric acid: 2NaBr + 2H 2 SO 4 + MnO 2 = Na 2 SO 4 + MnSO 4 + 2H 2 O + Br 2 The electrolysis of bromides also yields bromine just as that of chlorides yields chlorine. The reaction by which a bromide is converted into a chloride when treated with free chlorine is an important one. The fact that the reaction takes place shows clearly that chlorine is a more active element than bromine. When the reaction takes place in solution it occurs between bromine ions and chlorine molecules and indicates that the latter has a greater tendency to exist as an ion than the former has. The equation for the reaction, 2Na + + 2Br~ + C1 2 = 2Na + + Br 2 + 2CP resembles in this respect that of the reaction between salts and metals, for example, the action of iron on a solution of a copper salt : Cu + f + 2Cr + Fe = Cu + 2C1~ + Fe + + In the latter case we have learned that the reaction is said to take place because iron has a greater tendency than has copper to be an ion its solution pressure is higher. We now see that elements that pass into solution as negative ions behave in an analogous manner. Chlorine has a greater solution pressure than bromine. The elements that form negative ions can be arranged in a series THE HALOGEN FAMILY 367 based on this property in the same way as the metallic elements (574). 402. Physical Properties. Bromine, atomic weight 79.92, is a reddish-brown liquid (sp. gr. 3.12 at 20), which boils at 59, and freezes to needle-shaped crystals at 7.3; it possesses a most disagreeable and pungent odor. One hundred cubic centimeters of water dissolve 4.3 grams of the liquid at 0, and about 3 grams at ordinary temperatures; the solution, called bromine-water, gives off the halogen freely in the form of a vapor, which has a reddish color. Bromine has a high vapor pressure at room- temperature; at 18 the pressure is 150 mm., and, as a conse- quence, the gas over liquid bromine contained in a bottle at this temperature is approximately one-fifth bromine and four-fifths air. 403. Chemical Properties. Bromine affects markedly the tis- sues of which the body is made up. It has a very irritating effect on mucous membrane, and the inhalation of its vapor produces most disagreeable effects. When left in contact with the skin bromine causes painful sores which heal slowly. In chemical conduct bromine closely resembles chlorine; it combines directly with metals and many non-metals to form bromides, the properties of which are similar to those of the cor- responding chlorides. It forms at low temperatures a hydrate with water, Br2,10H2O, which is less stable than the hydrate of chlorine. 404. Uses. Bromine could be used for many of the purposes to which chlorine is put, but this is not done unless there are some particular advantages which come from its use. Bromine is employed in the preparation of bromides, some of which are exten- sively used. Potassium bromide is used in medicine as a sedative, and silver bromide is the most important substance in the sensitive film of a photographic plate. Free bromine is used in certain metallurgical processes and in the preparation of several important dyes and other organic substances. HYDROBROMIC ACID 405. Preparation. Hydrobromic acid, which is a colorless gas, is not made in a way analogous to that used for preparing hydro- chloric acid, namely, by treating a bromide with concentrated 368 INORGANIC CHEMISTRY FOR COLLEGES sulphuric acid. A similar reaction takes place in this case, but the acid formed is in part oxidized by sulphuric acid to bromine and water. It can be prepared conveniently by reactions entirely analogous to those into which chlorine enters, but which are not used for the preparation of hydrochloric acid. Bromine reacts with hydrocarbons and forms substitution-products and the acid. Benzene, CeHe, can be used for this purpose : C 6 H 6 + Br 2 = C 6 H 5 Br + HBr A similar reaction takes place smoothly in the cold when anthra- cene, CnHio, a product derived from coal-tar, is used. Hydrobromic acid is conveniently prepared by the action of water on the bromides of phosphorus. When the acid is made in this way bromine is first allowed to drop slowly on red phosphorus, which is converted into phosphorus pentabromide; water is next added, drop by drop, and as a result hydrobromic acid is formed: PBr 3 + 3H 2 O = HsP0 3 + 3HBr 406. Properties. Hydrobromic acid is a colorless gas, which fumes in the air, and has a sharp odor. By reduction in tempera- ture it can be converted into a liquid, which boils at 69. The gas is very soluble in water, 600 volumes dissolving in 1 volume of the latter at 0. The methods described above are used when hydrobromic acid is required in the form of a gas; if an aqueous solution is desired the reaction between a bromide and sulphuric acid may be used. This is possible because the oxidizing power of sulphuric acid is markedly reduced in the presence of water. It will be recalled that this fact is an example of a general truth; ions, in general, do not act as oxidizing agents, and when sulphuric acid is dissolved in water the molecules are largely ionized. When it is desired to make an aqueous solution of hydrobromic acid, potassium or sodium bromide, water, and sulphuric acid, in the correct propor- tions, are brought together and the mixture is distilled. The acid and water form a constant-boiling mixture (142) which con- tains approximately 48 per cent of HBr, has the specific gravity 1.49, and boils at 126 at 760 mm. pressure. In making the acid in this way the materials used are mixed in the proportion of 1 mol of the salt, 1 of sulphuric acid, and 5 of water, and distilled. THE HALOGEN FAMILY 369 407. It has been pointed out that bromine is a less active ele- ment than chlorine; its compound with hydrogen is, thus, more readily decomposed than hydrochloric acid. This is clearly seen in its behavior with oxidizing agents. Concentrated sulphuric acid, which is not sufficiently active as an oxidizing agent to affect hydrochloric acid, oxidizes hydrobromic acid. The equation for the reaction may be analyzed as follows: H 2 S0 4 = H 2 + SO 2 + [O] 2HBr + [O] = H 2 O + Br 2 H 2 SO 4 + 2HBr => 2H 2 O + SO 2 + Br 2 408. Properties of Bromides. The bromides of the common metallic elements resemble in physical properties the corresponding chlorides. They are all more or less easily soluble in water except those of silver, lead, and mercury (mercurous bromide, HgBr). They are less stable towards heat than the chlorides. The bromides of the acid-forming elements also resemble the chlorides of these elements. They have higher boiling-points, and are decomposed by water with the formation of hydrobromic acid and the acid derived from the other electro-negative element which they contain. 409. Test for Bromides. When a solution of silver nitrate is added to a solution of a bromide, silver bromide is precipitated in the form of a cloud or as a curdy solid. The solid, which has a light yellow tint, is insoluble in nitric acid. The reaction, which resembles closely that which takes place in the case of a chloride, is as follows : NaBr + AgNO 3 = AgBr + NaNO 3 Silver chloride is readily soluble in dilute ammonia, whereas silver bromide is difficultly soluble. If an unknown substance behaves with silver nitrate in a way to indicate that it is a bromide, another portion of its solution is treated, drop by drop, with chlorine-water, and shaken with 2 or 3 c.c. of carbon disulphide or chloroform. The bromine set free in the reaction, 2NaBr + C1 2 = 2NaCl + Br 2 , 370 INORGANIC CHEMISTRY FOR COLLEGES passes into the carbon disulphide and becomes evident on account of its red color. This behavior is shown by all substances which yield bromine ions when dissolved in water. In adding chlorine- water to a solution of a bromide in making the test, it is necessary to avoid the use of more chlorine than is required to liberate the bromine an excess of the reagent; in the presence of water, chlorine converts bromine into a colorless compound (431). IODINE 410. Occurrence and Discovery. Iodine occurs in the com- bined state in sea-water in very small quantities. It is present in certain aquatic plants and animals; and the former have been for many years a source of the element. In 1811 Courtois, a French manufacturer of soap, carried out some experiments with one of the products used in making soap, namely, the ashes obtained by burning kelp, which is a variety of sea-weed, that was used as a source of the potash required. He treated the ashes with concentrated sulphuric acid and discovered, as a result, that a violet-colored vapor was given off. Later, in 1814. Gay-Lussac published a full account of the substance discovered by Courtois, and proved that it was an element. The name selected for the newly discovered element, iodine, was derived from the Greek word meaning violet. Iodine is still obtained from the ashes of kelp, which contain from 0.5 to 1.5 per cent of a mixture of sodium and potassium iodides. The chief source at present is crude Chile saltpeter, NaNO 3 , which contains about 0.2 per cent of sodium iodate, NaIO 3 . Iodine occurs in the human body in traces, its presence in the thyroid gland being of great significance. In the disease known as goiter this gland is ill developed; medical research has shown that if a substance, called iodothyrine, which is obtained from the thyroid of the sheep, is injected, the human gland develops and the disease which accompanies its inactivity disappears to a large extent. 411. Preparation. Iodine is prepared from iodides by reac- tions entirely analogous to those used in the preparation of bromine. In the industrial preparation from the ashes of kelp, the latter are THE HALOGEN FAMILY 371 treated with manganese dioxide and sulphuric acid. The mixture is heated to the temperature at which the halogen sublimes, and the violet-colored vapor formed condenses to black crystals. In the laboratory potassium iodide is usually used in preparing iodine in this way. The iodine can also be liberated by treating the iodide with chlorine, and this process is used industrially. In the man- ufacture of iodine from crude Chile saltpeter the material is dis- solved in water and treated with a mixture of sodium sulphite and sodium bisulphite. To obtain iodine from sodium iodate, NalOa, it is necessary to use a reducing agent to remove the oxygen from the compound; the sulphites serve this purpose. The iodine which precipitates as a black solid, is filtered off, dried, and sublimed. In purifying iodine it is mixed with potassium iodide and resublimed. If any chlorine is present in the iodine, as iodine chloride, IC1, the latter reacts with the iodide as follows : KI + IC1 = KC1 + I 2 412. Physical Properties. Iodine, atomic weight 126.92, is obtained in the form of black crystalline plates when its vapor is condensed. It has the specific gravity 4.93, melts at 114, and boils at 184. The vapor of iodine has a violet color and is irri- tating when inhaled. A saturated aqueous solution of iodine contains 0.32 gram of the halogen per liter and has a light brown color; it is more soluble in alcohol, ether, chloroform, and carbon disulphide. Solutions of iodine in solvents which contain oxygen, such as water, ether, and alcohol, are brown in color, whereas solutions in other solvents are violet. The brown color is thought to be due to the presence in the first class of solvents of more or less stable addition-products of the solvent and the dissolved iodine. In the case of solvents which contain no oxygen the halo- gen has the same color that it would have if it were in the form of vapor. Iodine is very soluble in solutions of potassium iodide and other iodides, the amount dissolving being determined by the concentration of the iodide present. Such solutions are brown in color, and it has been shown that a compound is present which is formed as the result of the direct addition of the iodide and iodine: KI + I 2 <= KI 3 372 INORGANIC CHEMISTRY FOR COLLEGES The reaction is a reversible one, for if a solution of potassium iodide is saturated with iodine, and then diluted, a part of the latter is precipitated. A solution of iodine in potassium iodide is used in many processes in analytical chemistry. The indicator used in such processes, which are classed together under the name iodimetry, is a dilute solution of starch. When a few cubic centimeters of such a solution are added to a solution containing free iodine, a characteristic blue color is produced. It is probable that the color is formed as the result of the physical adsorption of the halogen by the starch. The reaction which serves to show the presence of exceedingly small amounts of iodine is also used to test for the presence of starch; in this case a dilute solution of iodine is added to the material tested. 413. Chemical Properties. The family relationship between chlorine, bromine, and iodine is clearly shown in the chemical properties of the latter. It is much less active than the other halogens, and any differences in behavior can be traced to this cause. Iodine forms iodides with metals and certain non-metals, but these are less stable than the corresponding bromides. It combines with hydrogen very slowly at elevated temperatures to form hydrogen iodide. The affinity of iodine for hydrogen is so small that under ordinary conditions it does not react with hydro- carbons in the way that chlorine and bromine do. The water-solution of iodine shows some of the properties of similar solutions of bromine and chlorine, but the proportion of hypoiodous acid formed is excessively small, and its presence becomes evident only when something is brought into contact with it with which it will react. For example, a solution of sul- phur dioxide in water (sulphurous acid) will decolorize a solution of iodine: I 2 + H 2 O <= [HIO] + HI [HIO] + H 2 SO 3 <=* HI + H 2 SO 4 I 2 + H 2 O + H 2 SO 3 <=* 2HI + H 2 SO 4 The reaction is a reversible one and if it is to run to completion from left to right, the solution must be dilute and an excess of iodine used. The reaction is commonly employed in determining quantitatively the amount of sulphurous acid present in a given THE HALOGEN FAMILY 373 solution. The amount of iodine added can be calculated if a definite volume of a solution of iodine of known strength is used; and the amount left over can be determined by titrating the solu- tion with a solution of sodium thiosulphate. The reaction which takes place is as follows : I 2 + 2Na 2 S 2 O 3 = 2NaI + Na 2 S 4 6 The difference between the amounts of iodine used in the two opera- tions is the quantity which was required to oxidize the sulphurous acid, and is, therefore, a measure of the amount of the latter. Many substances are oxidized in aqueous solution by iodine, and the process outlined above is an important one in quantitative analysis. Iodine in the presence of water reacts with other substances that are easily oxidized. For example, hydrogen sulphide and iodine form hydriodic acid and sulphur: H 2 S + I 2 = 2HI + S The reaction expressed by the above equation may be utilized to prepare an aqueous solution of hydriodic acid. A solution of hydrobromic acid can be made by an analogous reaction. For this purpose the halogen is covered with the amount of water desired as the solvent for the acid formed, and a stream of hydrogen sulphide gas is passed into the solution. When the reaction is complete the sulphur is filtered off and the liquid is distilled. 414. Uses. Iodine affects the skin and mucous membrane, but its action is not so marked as that of bromine. It reduces swellings and hardens the skin. It was much used in the army to relieve the pain which results from abrasions of the feet induced by long marches. The affected parts were treated with an alco- holic solution (tincture) of iodine. Iodine, as well as a number of organic compounds which contain the element, has marked antiseptic properties; iodoform, CHIs, which is extensively used in surgery for this reason, is manufac- tured from alcohol, sodium hydroxide, and iodine. Sodium, potassium, and rubidium iodides are used in medicine. HYDRIODIC ACID 415. Preparation. Hydriodic acid, which is a colorless gas, is not prepared by the action of sulphuric acid on an iodide; even in 374 INORGANIC CHEMISTRY FOR COLLEGES aqueous solution the acid oxidizes the hydriodic acid formed and iodine is set free. It cannot be prepared by the action of iodine on a hydrocarbon in a way analogous to that by which hydrogen bromide can be made. The gas can be prepared by heating sodium iodide with concentrated phosphoric acid, HaPCU. In this case a double decomposition takes place; as an excess of the acid is used an acid phosphate is formed : Nal + H 3 PO 4 = NaH 2 PO 4 + HI We shall learn later that phosphoric acid is not an oxidizing agent; as a consequence, it does not react with hydrogen iodide. The method commonly employed to make the gas is to treat with water, phosphorus iodide prepared by the action of iodine on phosphorus : PI 3 + 3H 2 = P(OH) 3 + SHI 416. Physical Properties. Hydriodic acid is a heavy, colorless gas which fumes in the air. It can be condensed to a liquid which boils at -34 and freezes at -51. One volume of water at 10 dissolves 425 volumes of the gas, and the resulting solution con- tains 70 per cent of hydrogen iodide. It forms a constant-boiling mixture with water, which has the specific gravity 1.7, contains 57 per cent of hydrogen iodide, and boils at 127 at 760 mm. pres- sure. 417. Chemical Properties. Hydriodic acid resembles in its general behavior hydrochloric acid and hydrobromic acid; the differences observed are due to the fact that the latter are much the more stable substances. Hydrogen iodide begins to decompose slowly into hydrogen and iodine at 180; at higher temperatures the reaction is a rapid one. This fact can be easily shown by pouring from a large jar a quantity of the gas on a Bunsen flame. As the heavy gas falls on the flame it becomes heated, is decom- posed, and a violet cloud of iodine vapor is seen. If gaseous hydriodic acid is mixed with chlorine a reaction takes place with the evolution of light: 2HI + C1 2 = 2HC1 + I 2 In aqueous solution both the acid and its salts are decomposed by chlorine in this way; bromine acts in a similar manner. THE HALOGEN FAMILY 375 Hydriodic acid is a strong acid, and resembles closely hydro- chloric acid. Its solution is unstable in the presence of the air on account of the fact that oxygen reacts with it slowly at ordinary temperatures : 4HI + O 2 = 2H 2 O + 2I 2 The reaction takes place more rapidly in the presence of sunlight, and to decrease the rate of the decomposition the solution is usually kept in an opaque container or a bottle of brown glass. Oxidizing agents liberate iodine from hydriodic acid. Its behavior with sulphuric acid has been noted. When solid potassium iodide is treated with concentrated sulphuric acid, the presence of sulphur dioxide, sulphur, and hydrogen sulphide can be noted ; there is also present some hydrogen iodide which escapes decomposition. A dilute solution of hydriodic acid is used in medicine to pro- mote certain secretions. The acid is used in organic chemistry to prepare iodides and to test for the presence of certain important groupings of elements; it is one of the most active reducing agents and is used for this purpose when other agents fail. 418. Test for Iodides. The reactions upon which the test for iodides is based are analogous to those used in the case of bromides. When silver nitrate is added to a solution of an iodide a precipitate of silver iodide is formed, which is insoluble in nitric acid and ammonia, and has a greenish-yellow color. If the behavior of a salt with silver nitrate indicates that it is an iodide, to a second portion of the solution is added a few cubic centimeters of carbon disul- phide; chlorine-water is next added drop by drop, and the solution shaken. If an iodide is present iodine is set free, dissolves in the carbon disulphide, and is evident as the result of the violet color of its solution. If a solution contains both a bromide and an iodide, the iodine is first set free as the chlorine-water is slowly added; when this has reacted with the added chlorine to form a colorless com- pound (431), the bromine is liberated and imparts to the carbon disulphide its characteristic red color. FLUORINE 419. The element which has the smallest atomic weight in the family of the halogens is fluorine, and on this account it might 376 INORGANIC CHEMISTRY FOR COLLEGES have been considered first in a discussion of the group. Its con- sideration has been delayed, however, for a good reason. Fluorine is an exceedingly active element and, consequently, it enters into reactions not exhibited by chlorine, and its compounds are so stable that many of them do not show the behavior exhibited by the chlorides of similar composition. Chlorine, on the other hand, as we have seen, does not differ markedly in activity from bromine and a consideration of the latter after the chemistry of chlorine has been mastered gives an excellent opportunity to contrast the behavior of the two elements and their compounds, and to bring out relationships indicated by the periodic law. We have also seen that while iodine resembles chlorine and bromine in many respects, its lack of activity results in a behavior which is char- acteristic. We are now in a position to study the chemistry of an element similar to, but much more active than chlorine. In this case we shall see that fluorine and its compounds will not enter into certain reactions analogous to those shown by the other halogens, and new types of reactions will appear. In any chemical family the element having the lowest atomic weight shows, in general, under a given set of conditions, the greatest activity. The differ- ence between its behavior and that of the next member of the group is much greater than that between succeeding members. For this reason the members of the first series of elements in the peri- odic classification are not as typical of the families in which they occur as the members of the second series are. In general, the groups of elements which show the most striking relationships are composed of members from the second to the tenth series. 420. Occurrence and Discovery. Fluorine occurs in small quantities rather widely distributed. It is present in several minerals, the most important of which are fluorspar or fluorite, CaF 2 , cryolite, AlF 3 ,3NaF, and apatite, CaF 2 ,3Ca 3 (PO4) 2 . Fluorine also occurs in traces in the human body, and is present in appreciable amounts in the enamel of the teeth. It has been suggested that when the teeth of growing children do not develop normally one cause may be a lack in the food of sufficient com- pounds containing fluorine. The body contains a large number of elements in small quantities and we must obtain these from food. Under most circumstances an adequate supply of these elements is THE HALOGEN FAMILY 377 obtained, but when the diet is restricted to a few substances only there may result a deficiency in one of them. The alchemists were familiar with the fact that when fluorspar is heated with concentrated sulphuric acid a colorless gas is formed, which has the unusual property of etching and dissolving glass. Gay-Lussac studied the reaction in 1807 and isolated the gas which was called hydrofluoric acid, the name indicating its preparation from fluorspar. The mineral had been previously given this name because it can be easily melted the first two syl- lables in fluorspar being derived from the Latin word to flow, and the last being the general name for minerals. For over seventy-five years chemists attempted in vain to isolate from fluorides the undiscovered element which was thought to be present. All attempts were unsuccessful on account of the great activity of the element, which was called fluorine. Finally, Moissan in 1886 came to the conclusion that the electric current must be used to effect the decomposition of fluorides, and that this must be carried out in the absence of water. When potassium fluoride was dissolved in liquid hydrogen fluoride and subjected to the action of an electric current, a light yellow gas was formed, which proved to be the element sought. 421. Preparation. No methods similar to those employed to prepare chlorine can be used for the isolation of fluorine from its compounds. Oxidizing agents do not convert hydrofluoric acid into fluorine and water, because fluorine reacts quantitatively with water to form hydrofluoric acid and oxygen. For this reason, also, electrolysis of an aqueous solution of the acid does not yield the free halogen. Fluorine is prepared in the way used by Moissan. The appa- ratus in which the electrolysis is carried out is made of copper. Although fluorine attacks the metal, it soon becomes coated with a layer of copper fluoride, which does not dissolve in the hydro- fluoric acid used as the solvent in the electrolysis and serves, there- fore, as a protective coating for the metal. The electrodes used are made from an alloy of platinum and iridium as all other sub- stances that have been tried are attacked by the nascent fluorine. The reaction is carried out at a low temperature in order to reduce the activity of the halogen- 378 INORGANIC CHEMISTRY FOR COLLEGES 422. Physical Properties. Fluorine, F 2 , atomic weight, 19, is a gas with a pale yellow color; it has been converted into a liquid which boils at -186 and freezes at -223. 423. Chemical Properties. Fluorine reacts with all elements except chlorine, oxygen, nitrogen, and the inert gases of the helium family. When perfectly dry it does not attack silicon dioxide or glass, but in the presence of a trace of moisture some hydrofluoric acid is formed, and, as a result, the silicon dioxide is dissolved and the glass, which contains this oxide, disintegrated. It is for this reason that glass vessels cannot be used in preparing the element. The noble metals are converted into fluorides very slowly by the gas. Fluorine reacts with water instantaneously; when the gas is passed into water the chief products of the reaction are oxygen and hydrofluoric acid, which we shall see later has the formula H2F2: 2F 2 + 2H 2 O = 2H 2 F 2 + O 2 If water-vapor is brought into contact with fluorine, ozone is formed. On account of the difficulty attending the preparation of fluorine and the great activity of the gas, it is not used for any purpose. The chief interest in fluorine centers in its remarkable chemical activity and the relation of its properties to those of the other members of the halogen family. HYDROFLUORIC ACID 424. Preparation. The formula of hydrofluoric acid, H 2 F 2 , indicates that it is a dibasic acid. Like acids of this type it forms acid salts, one of which has the formula KHF 2 . Many such acid salts when heated are broken down into the neutral salt and the free acid. Advantage is taken of this behavior of the acid fluoride to prepare the anhydrous acid: 2KHF 2 = 2KF + H 2 F 2 Hydrofluoric acid is generally used in solution, and is com- monly prepared by heating powdered calcium fluoride with con- centrated sulphuric acid in a retort of lead, which is but slightly attacked by the gas; the latter is absorbed in water contained in a bottle made of hard rubber or wax. The reaction which takes place is analogous to that between a chloride and sulphuric acid. THE HALOGEN FAMILY 379 425. Physical Properties. Hydrofluoric acid is a very volatile, colorless liquid which boils at 19.4. It forms a constant-boiling mixture with water, which boils at 111 at 750 mm. pressure and contains 43.2 per cent of the acid. The molecular weight of hydrogen fluoride, as determined by the density of the gas, varies markedly with the temperature; the volume of the gas at 26 which occupies 22.4 liters when reduced to and 760 mm. weighs 51 grams; a similar volume of the gas at a temperature above 90 weighs 20 grams. These facts indicate that above 90 the molecular weight of the gas is 20, and its formula HF. Below 90 more complex molecules are present; the molecular weight of a compound having the formula H2F2 is 40 and HaFa is 60. The conclusion is drawn, therefore, that in hydrogen fluoride these three types of molecules are present in equilibrium: H 3 F 3 + H 2 F 2 <= HF. The formula H 2 F 2 is usually assigned to hydrofluoric acid because the molecular weight of the acid determined by the freezing-point method corresponds closely to this formula. When two or more molecules of the same kind unite to form a single molecule that reverts more or less completely, with slight change in the conditions, to the molecules from which it was formed, the phenomenon is called association. The vapor of hydrofluoric acid from a temperature just above the boiling-point of the liquid to about 90 is associated because it is made up of the molecules HF, H 2 F 2 , and HaFa in equilibrium. It will be recalled that the same phenomenon was seen in the case of nitrogen tetroxide the vapor of which contains two kinds of molecules in equilibrium N 2 C>4 <= 2N0 2 . Associated liquids are also known; it is highly probable that water contains different kinds of molecules in equilibrium HeOs + H4O 2 + H 2 O. Many liquids are known the properties of which can best be explained on the assumption that they are associated. The formulas of these liquids are written in such a way as to indicate their relation to the formulas of the simple substances of which they are made up. For example, in the case of water the associated molecules are usually represented by the formulas (H 2 O)s and (H 2 O) 2 and not as HeOs and H4O 2 as given above. For the sake of simplicity such formulas are used only when the fact that the liquid is more or less associated, is an important factor in the reaction under consideration. 380 INORGANIC CHEMISTRY FOR COLLEGES 426. Chemical Properties and Uses. When hydrofluoric ?cid is dissolved in water, the acid in solution is much weaker than hydrochloric acid. At 18 in one-tenth normal solution, the latter is dissociated into its ions to the extent of about 92 per cent; under the same conditions hydrofluoric acid is 0.15 per cent dis- sociated. This fact is a result, in all probability, of the great activity of fluorine, but the exact reason is not apparent. The fact that the acid exists in the associated form is also traceable to the same cause. This view is expressed by considering the complex mole- cule as made up as indicated in the following graphic formula: H F = F H. From this point of view we would not expect that the ionization of this complex molecule would necessarily bear any relation to that of the smaller molecule of hydrochloric acid, HC1. Hydrofluoric acid is a sufficiently active acid to react with the more active metals. It forms fluorides with oxides and hydroxides of the metals. It differs from hydrochloric acid in being dibasic and, accordingly, forms acid salts of which KHF2 is an example. A very characteristic property of hydrofluoric acid is shown in its action on silicon dioxide, which occurs in an impure form as sand. The reaction is represented by the following equation: SiO 2 + 2H 2 F 2 = SiF 4 + 2H 2 O Silicon fluoride is a gas, and, as a result, when pure silicon dioxide is treated with hydrogen fluoride in the form of a gas or in solution, the oxide is converted into substances which are volatile. This reaction is utilized in quantitive chemical analysis in determining the amount of silicon present in substances containing the element. The latter is separated as silicon dioxide in a more or less pure condition, and is heated until its weight remains constant; the material is then treated with a solution of hydrofluoric acid, and heated again to constant weight; the loss in weight equals the weight of the pure silicon dioxide which was present, and from it can be calculated the amount of silicon in the compound analyzed. 427. Salts of the acids containing oxygen may be considered as made up of oxides of metals and of the anhydrides of the acids. Considered in this way the formula CaSiOs may be written as THE HALOGEN FAMILY 381 follows: CaO,SiO 2 . When the salt is treated with hydrofluoric acid both oxides are converted into fluorides: CaO,SiO 2 + 3H 2 F 2 = CaF 2 + SiF 4 + 3H 2 O Common glass is a mixture of sodium silicate, NaoSiOa, calcium silicate and silicon dioxide. When it is treated with hydrofluoric acid, the silicon passes off as a gas in the form of the tetrafluoride, SiF-i, and the calcium fluoride and sodium fluoride are left behind in the form of a powder. As the result of the reaction the glass is eaten away, or etched, wherever hydrofluoric acid comes in con- tact with it. The reaction is used to etch designs upon glass. To do this the glass is covered with a thin layer of wax and the design is drawn on the latter with a sharp point which cuts through the wax to the glass. When exposed to hydrogen fluoride the uncovered part of the glass is attacked by the acid, and when the wax has been removed, the design is clearly seen as the result of the fact that the etched surface is rough and semi-transparent like ground glass. It is in this way that the markings are put on chemical thermometers, burettes, and other measuring instru- ments made of glass. 428. Properties of Fluorides. The fluorides differ markedly from the chlorides, bromides, and iodides in solubility and other physical properties. Silver fluoride is a very soluble salt; its molar solubility is about two and one-half times that of common salt. On the other hand, calcium fluoride is very difficultly soluble, whereas the other halides of the metal dissolve readily in water. The fluorides of the metals are very stable towards heat; those of the acid-forming elements react with water in a way similar to that in which the analogous compounds of the other halogens react. The test for fluorides is based on the action of hydrofluoric acid on glass. The substance to be tested is ground to a fine powder and mixed with concentrated sulphuric acid in a lead dish. A glass plate coated with a thin layer of wax or paraffin, in which a mark has been scratched, is placed over the dish. After a few minutes the wax is removed from the glass, and if a fluoride was present the mark will appear etched into the glass. 382 INORGANIC CHEMISTRY FOR COLLEGES COMPOUNDS OF THE HALOGENS WITH OXYGEN AND WITH HYDROGEN AND OXYGEN 429. Fluorine forms no compound with oxygen, although chlorine and iodine do form such compounds. Acids which con- tain chlorine, bromine, or iodine in combination with hydrogen and oxygen are known; they form well-characterized salts, some of which have important uses. Attempts to prepare analogous compounds containing fluorine have been unsuccessful. 430. Chlorine Monoxide. When chlorine is dissolved in water a reaction takes place as the result of which hypochlorous acid and hydrochloric acid are formed (125) ; C1 2 + H 2 O = HOC1 + HC1 If sodium hydroxide is added to the solution, or if chlorine is passed into a solution of the base, the salts of the two acids are formed. Sodium hypochlorite is an unstable substance. When it is treated with sulphuric acid, the hypochlorous acid set free de- composes into its anhydride and water: 2NaOCl +~H 2 SO 4 = Na 2 SO 4 + 2HOC1 2HOC1<=>H 2 + C1 2 Chlorine monoxide, C1 2 0, is a brownish-yellow gas (boiling- point 5) which reacts with water as indicated above. It is best prepared by the action of chlorine on freshly precipitated mercuric oxide, which has been carefully dried: HgO + 2C1 2 = HgCl 2 + C1 2 The oxide is a very unstable substance, and explodes readily. 431. Hypochlorous Acid. We have just learned that this acid is an unstable substance and decomposes into its anhydride and water when liberated from its salts. It can, however, exist in dilute aqueous solution. It decomposes when exposed to sun- light (125), 2HOC1 = 2HC1 + 2 , and when brought into contact with substances that can be oxidized, and for this reason it is used as a bleaching agent. Hypochlorous acid is very slightly dissociated, a fact which is utilized in preparing dilute solutions of the acid. This is accom- THE HALOGEN FAMILY 383 plished by passing chlorine into a dilute solution of the salt of a weak acid which is more highly dissociated than hypochlorous acid. Sodium carbonate can be used, for example. The reactions involved are as follows: 2C1 2 + H 2 O + [H 2 O] <= 2HOC1 + [2HC1] Na 2 CO 3 + [2HC1] = 2NaCl + [H 2 O] + C0 2 Na 2 CO 3 + 2C1 2 + H 2 O = 2HOC1 + 2NaCl + C0 2 The sodium carbonate reacts with the hydrochloric acid formed to produce sodium chloride, but does not appreciably affect the weak hypochlorous acid. The solution so produced can be used directly, if the presence of sodium chloride does not interfere with its use, or it can be distilled; if it is dilute the distillate obtained is a dilute solution of hypochlorous acid. Hypochlorous acid can be prepared by treating bleaching powder with a weak acid. When bleaching powder (639) is dis- solved in water there are two salts present, calcium hypochlorite, Ca(OCl) 2 , and calcium chloride, CaCl 2 . If to such a solution an acid is added stronger than hypochlorous acid and weaker than hydrochloric acid, the former will be set free and the latter will not; carbonic acid (CO 2 + H 2 O), can be used for this purpose: Ca ++ + 2(001)" + 2H + + C0 3 "~ = CaCO 3 + 2HOC1 Hypochlorous acid is an active oxidizing agent. A number of reactions have already been considered in which the acid acts in this way (126). It oxidizes iodine in aqueous solution to iodic acid: 5HOC1 + I 2 + H 2 O <=> 2HIO 3 + 5HC1 A similar reaction takes place with bromine, and bromic acid, HBrO 3 , is formed. It is for this reason that chlorine-water is added cautiously in testing for bromides and iodides; an excess of the former would convert the bromine or iodine liberated into bromic acid or iodic acid, both of which are colorless, and, as a result, the presence of the halogens would be overlooked. 432. Hypochlorites. Most of the salts of hypochlorous acid are unstable, and when they are to be employed for any purpose 384 INORGANIC CHEMISTRY FOR COLLEGES they are made in solution and used in this form. They react with hydrochloric acid and form chlorine: NaOCl + 2HC1 = NaCl + C1 2 + H 2 O The hypochlorite of sodium is prepared by passing chlorine into a cold, dilute solution of sodium hydroxide: C1 2 + [H 2 O] <=* [HC1] + [HOC1] NaOH + [HC1] = NaCl + [H 2 O] NaOH + [HOC1] = NaOCl + H 2 O 2NaOH + C1 2 = NaCl + NaOCl + H 2 O 433. Chloric Acid and Chlorates. When an aqueous solution of sodium hypochlorite is heated, sodium chlorate is formed. The hypochlorites of other metals behave in a similar way. The chlorate commonly prepared and used is the potassium salt; on account of the fact that its solubility is relatively small in cold water, it can be separated from the potassium chloride formed along with it. The salt can be made by passing chlorine into a strong hot solution of potassium hydroxide. The reactions which take place can be analyzed as follows : 3C1 2 + [3H 2 0] = [3HC1] + [3HC10] 3KOH + [3HC1] = 3KC1 + 3H 2 O 3KOH + [3HC10] = [3KC1O] + [3H 2 O] [2KC1O] = 2KC1 + [20] [KC10] + [20] = KC10 3 3C1 2 + 6KOH = KC1O 3 + 5KC1+ 3H 2 O After completion of the reaction the potassium chlorate is obtained in the form of crystals when the solution is cooled. This method of preparing potassium chlorate is an expensive one on account of the fact that only one-sixth of the chlorine and one-sixth of the potassium hydroxide are converted into the desired salt; the rest is obtained in the form of potassium chloride, a salt that occurs in nature and is, accordingly, less expensive than potassium hydroxide. Potassium chlorate is made electrolytically on the large scale from potassium chloride. When a current is passed through a THE HALOGEN FAMILY 385 solution of the chloride, oxidation takes place at the anode and potassium chlorate is formed; hydrogen is set free at the cathode. 434. Chloric acid, HC1O 3 , is known only in solution; when an attempt is made to concentrate the latter it decomposes when about 40 per cent of the acid is present. It also decomposes when heated above 40. For these reasons chloric acid cannot be ob- tained from its salts by the use of the method employed in the case of acids which boil without decomposition. It is prepared by double decomposition in aqueous solution from a chlorate and an acid, the two substances being selected so that the metal in the salt and the radical in the acid are those of an insoluble salt; the latter precipitates and the desired acid remains in solution. For example, barium chlorate and sulphuric acid could be used: Ba(ClO 3 ) 2 + H 2 SO 4 - BaSO 4 + 2HC1O 3 If potassium chlorate is treated with concentrated sulphuric acid, the chloric acid set free decomposes spontaneously with explosive violence; among the products formed are chlorine dioxide, C1O2, and oxygen. Chloric acid is an active oxidizing agent; a strong aqueous solution of the acid will ignite paper. 435. When potassium chlorate is heated to a high temperature it loses its oxygen: 2KC10 3 = 2KC1 + 30 2 At lower temperatures a reaction takes place which resembles somewhat the change that occurs when hypochlorites are heated. In the case of the chlorate reduction takes place simultaneously with oxidation and a perchlorate is formed: 4KC1O 3 = KC1 + 3KC1O 4 Chlorates are active oxidizing agents. Mixtures made up of potassium chlorate and readily oxidizable substances, such as red phosphorus and sulphur, explode violently when struck. Such mixtures diluted with sand or fine pebbles, are used in making toy torpedoes for use in celebrating Independence Day. 436. Perchloric Acid and Perchlorates. Potassium perchlorate is about one-twentieth as soluble as potassium chloride and can be readily separated by crystallization from the chloride formed 386 INORGANIC CHEMISTRY FOR COLLEGES when potassium chlorate is cautiously heated. The equation for the reaction by which potassium perchlorate is formed has been given (435) . Perchloric acid is a colorless liquid which can be heated to about 90 without change. On account of this fact and the fact that it boils below this temperature when the pressure is reduced, the acid can be prepared by treating potassium per- chlorate with concentrated sulphuric acid and distilling the acid under reduced pressure. The process has to be carried out with great care and the pressure must be as low as possible in order that distillation may take place at a low temperature. Per- chloric acid has been recently prepared commercially by the electrolytic oxidation of hydrochloric acid. In order to prevent decomposition, perchloric acid is usually kept in solution (70 per cent or less). The acid is an oxidizing agent, but not so active as the other oxygen acids of chlorine. When it decomposes, chlorine dioxide is among the products formed. When perchloric acid is treated at a low temperature with phosphorus pentoxide, which is a powerful dehydrating agent, the anhydride of the acid is formed, and the water removed from tho the acid unites with the pentoxide to form metaphosphoric acid : 2HC10 4 + P 2 O 5 = C1 2 7 + 2HP0 3 Perchloric anhydride is obtained from the mixture by careful distillation; it is a colorless liquid, which boils at 82, and explodes when struck or heated above its boiling-point. The salts of perchloric acid are more stable than the chlorates, but they lose their oxygen when heated and give it up readily to substances that can be oxidized; they are used in the preparation of various kinds of fireworks. On account of the fact that potassium perchlorate is a difficultly soluble salt, perchloric acid is used in the quantitative determina- tion of potassium. 437. Chlorine Dioxide. When chloric acid and perchloric acid decompose, chlorine dioxide, C1O2, a yellow gas which con- denses to a liquid at 10, is among the products formed. It explodes violently, and a large amount of heat is evolved; it is a highly endothermic compound. THE HALOGEN FAMILY 387 Chlorine dioxide and water react and form chlorous and chloric acids: 2C1O 2 + H 2 O = HClOo + HC1O 3 The chlorous acid formed in this way exists only in solution. From the acid, salts called chlorites have been prepared. 438. Hypobromous Acid and Bromic Acid. These acids and their salts are well known; the formulas and properties of the substances are similar in each case to those of the analogous chlorine compounds. The acids are active oxidizing agents, but less so than hypochlorous and chloric acid. No oxides of bromine have been isolated. 439. Oxygen Acids of Iodine. Hypoiodous acid, HIO, and its salts are known in solution, but they pass readily into iodic acid and iodates. They are active oxidizing agents, which resem- ble bromic acid and bromates in properties and modes of prepa- ration. Sodium iodate, NalOa, is an important salt, since it occurs in Chile saltpeter and is the chief source of the iodine of commerce. Iodic acid can be readily prepared by oxidizing iodine with strong nitric acid. It is a white solid which loses water at about 170 and passes into iodic anhydride: 2HIO 3 = H 2 At about 300 the anhydride begins to decompose into iodine and oxygen, but it does not explode. 440. Compounds of the Halogens with Themselves and with Other Elements. Two chlorides of iodine, the monochloride, IC1, and the trichloride, Ida, are formed as the result of the direct union of the halogens; the monochloride is a red crystalline sub- stance and the trichloride a yellow powder. They are decom- posed by water and by heat. Compounds of the formula IBr and IFs are said to exist. The halogens form compounds with other acid-forming ele- ments. Nitrogen chloride, NCls, is an oily liquid which is formed when a solution of ammonium chloride is treated with chlorine; it is highly explosive. No bromide of nitrogen is known. Nitro- gen iodide, to which the formula Nl3,NHs is assigned, is prepared by treating a strong aqueous solution of ammonia with iodine. It 388 INORGANIC CHEMISTRY FOR COLLEGES separates as a black insoluble precipitate, which decomposes vio- lently when touched in the dry condition. 441. The Halogen Family and the Periodic Law. The study of the properties of the halogens and their important compounds has brought out clearly the interesting relationships which exist among them. In general, the change in any property is progressive when we pass from one element to another in the order of increasing atomic weights. If the property is a physical one closely associated with weight, such as density, its value increases with increasing atomic weight ; if it is not so associated, the change may be in the reverse direction. For example, the solubilities of the halides of silver decrease as we pass from the fluoride to the iodide, but the solubilities of the halides of sodium increase as we pass in the same direction. We do not know the causes underlying solubility, but the fact is important that in most cases there is a progressive change in solubility as we pass from one compound to the next in a series of salts of analogous composition derived from the elements in a chemical family. The change in chemical properties is also progressive. For example, the activity of the halogens, as measured by the energy liberated when they unite with hydrogen, decreases with increasing atomic weight, and the stability of the compounds formed decreases in the same order. When we consider, however, the compounds containing oxygen, and oxygen and hydrogen, the order is reversed ; the most stable compounds are derived from iodine. As a result of these facts chlorine will drive out iodine from hydriodic acid, whereas iodine will convert chloric acid into iodic acid. The order of replacement is reversed in the second case on account of the fact that oxygen compounds containing iodine are more stable than the analogous chlorine compounds. According to the periodic table the elements in the seventh group have the valence 1 toward hydrogen and 7 toward oxygen. The valencies indicated in the table are the highest possible; in many cases compounds, the possibility of the existence of which is indicated by the law, have never been prepared. Perchloric anhydride, C^Oy, is the only oxide in which a halogen atom shows the valence 7. Valencies between 1 and 7 are shown in a num- ber of compounds. The graphic formulas of some of these are written as follows: THE HALOGEN FAMILY 389 C1 No H O Cl=0 Cl/ H-O-C1/ H-O-C1^=O CK ^O ^0 \0 Cl^O 442. The Electronic Theory of Valence. A very brief account has already been given of the recently advanced theory of the con- stitution of matter, which postulates that the atoms are made up of charges of electricity (398) . According to this view the unchange- able nucleus of the atom is composed of positive and negative charges of electricity, and outside of these, in what is called the outer sphere, there are a definite number of negative charges which come into play when the atom unites with another atom. These negative charges are called valence electrons, and the valence of an element is determined by the number of these electrons possessed by each atom of the element. The number of valence electrons varies from in the case of the inert elements in group of the periodic classification to 8 in the case of the elements in group 8, the number on any element being the same as the number of the group in which it occurs in the classification. According to this theory, when chemical union takes place between two elements, a valence electron passes from one element to the other, and there is established between the two a so-called electrical field of force, which holds them in combination. The element which loses the electron is the more positive of the two. When, for example, a sodium atom, which has 1 valence electron (group 1), unites with chlorine, which has 7 valence electrons (group 7) , the metallic atom loses its electron, which passes to the chlorine atom. This is expressed in the electronic formula for sodium chloride as follows: Na > Cl; the line ordinarily representing a valence bond is replaced by an arrow to indicate that in this case an electron has passed from the sodium atom to the chlorine atom. When the product of the reaction, sodium chloride, is dissolved in water, the latter separates the atoms from each other, and, as a 390 INORGANIC CHEMISTRY FOR COLLEGES consequence, the sodium which in its union with chlorine lost a negative charge, becomes positive, a sodium ion, and has the valence + 1 ; and the chlorine which has gained a negative charge becomes negative and has the valence 1. Losing a negative charge of electricity is the same as gaining a positive charge. On account of the fact that the compounds formed as the result of the union of two non-metallic elements do not dissociate into ions, the union between such elements is considered to be of a different nature. In this case the elements are said to share pairs of electrons between them. In chlorine monoxide the union is represented thus : Cl :O :C1. The two dots between chlorine and oxygen indicate that each element furnishes an electron which is held in common by the two atoms. There is not a sufficient dif- ference in chemical properties between the atoms to cause the passage of an electron from one element to the other (806). It will be seen by inspecting the periodic table that there is a remarkable fact in regard to the positive and negative valences of all the elements. In group 7 the elements form compounds of the type HX, in which the valence is 1 and of the type X2Oj in which the valence is + 7. In group 6 the valence values are 2 and + 6, in group 5 they are 3 and + 5, and in group 4,' 4 and + 4. The sum of the valencies disregarding the signs is 8 in each case. This fact is a help in remembering the valencies of elements; if, for example, we know that the highest valence of an element toward hydrogen is 3, then we know that its valence toward oxygen and other negative elements is 5. The facts given above lead to the conclusion that the largest number of valence electrons which can exist on an element is 8. For example, chlorine, which has 7 valence electrons, can receive but 1 more. EXERCISES 1. (a) At what temperature does bromine boil when the pressure on it is 150 mm.? (6) What weight of bromine is required to fill at 20 a bottle which holds 250 c.c.? 2. (a) Write an equation expressing the equilibrium between the sub- stances present in a solution of bromine in water. (6) Explain what happens if air is bubbled through the solution. 3. The average of a number of experiments led to the conclusion that the weight of 1 liter of hydrogen bromide reduced to and 760 mm. is 3.612 grams. Using 1.008 as the atomic weight of hydrogen and HBr as the formula of hydrogen bromide calculate the atomic weight of bromine. THE HALOGEN FAMILY 391 4. (a) State what would happen if finely divided silver is shaken with a solution of hydrobromic acid and bromine. (6) Devise a way to determine quantitatively the amount of each in the solution. 5. What (a) weight .and what (6) volume of the constant boiling aque- ous solution of hydrobromic acid are required to react with 10 grams of zinc? 6. The reaction between sulphuric acid and hydrobromic acid is a revers- ible one. Under what conditions could it be used to prepare a solution of hydrobromic acid from bromine? 7. What weight of iodine must be dissolved in 1000 c.c. of water to make (a) a 0.1N solution and (6) a 0.1 molar solution? What weight of Na 2 S 2 O 3 must be dissolved in 100 c.c. of water to make (c) a 0.1N solution and (d) a 0.1 molar solution? 8. Forty c.c. of a 0.05 molar solution of I 2 were added to 100 c.c. of an aqueous solution of SO 2 . The excess of I 2 required 15 c.c. of a 0.05 molar solution of Na 2 S 2 O 3 to react with it. Calculate (a) the weight of SO 2 in the solution, (6) percentage by weight of SO 2 present and (c) the volume of SO 2 required to prepare such a solution. 9. What weight of iodine is required to make 100 c.c. of the constant boiling solution of hydriodic acid? 10. Compare the action of (a) light and (6) heat on the reactions repre- sented by the following equations : C1 2 + H 2 O + 2HC1 + O 2 , and I 2 + H 2 O + 2HI + O 2 . 11. Devise a method of determining quantitatively the amount of a soluble fluoride in a solution. 12. What changes in color are observed when chlorine-water is added to a solution containing a bromide and an iodide? (6) Write equations for the reactions involved including those which lead to the decolorization of the solution. 13. Devise a way in which the chlorine set free could be determined quanti- tatively when a solution of bleaching powder is acidified. 14. How could you distinguish from each other solutions containing (a) sodium chloride, (6) sodium hypochlorite, (c) sodium chlorate, (d) sodium iodide, and (e) sodium bromide? 15. How could you distinguish the following from one another: (a) NaClO 3 , (6) NaIO 3 , (c) NaF? 16. Compare the behavior of C1O 2 and NO 2 when brought into contact with cold water. State the changes in valence of Cl and N which take place in the two cases. CHAPTER XXVII SELENIUM AND TELLURIUM 443. Selenium and tellurium are in the sixth group in the periodic classification , in which oxygen is the element with the lowest atomic weight; they form with sulphur a chemical family in which the relationships between the members of the family are as striking as in the case of the halogens. All elements in the family have a negative valence of 2 and can exhibit a positive valence up to 6. Selenium and tellurium form compounds anal- ogous in composition to [28, SO2, H^SOs, EfeSCU, SC12, etc. The elements are present in the earth in relatively small quantities only, and have not been used for many purposes. They and some of their compounds will be considered briefly. 444. Occurrence. Selenium occurs in the elementary condi- tions associated with free sulphur, and as a selenite along with sul- phides. It occurs in some varieties of pyrite, FeS2. It was dis- covered by Berzelius in 1817 in the flue dust separated from the gas produced by burning pyrite in the manufacture of sulphuric acid. Tellurium occurs largely in combination with copper, gold, or silver as tellurides, and is obtained as a by-product in the elec- trolytic refining of copper. The element was discovered in 1783 by Miiller von Reichenstein and was later named tellurium on account of the earthy appearance of the substance from which it was obtained (tellus, the earth). The name of selenium, discov- ered later, was derived from the Greek word signifying the moon. 445. Preparation and Physical Properties. In the case of both selenium and tellurium the usual method of preparation is the same, namely, the precipitation of the free element from a solution of the chloride by means of sulphur dioxide: SeCU + 2SO 2 + 4H 2 O = Se + 2H 2 SO 4 + 4HC1 392 SELENIUM AND TELLURIUM 393 The elements can be separated by utilizing the fact that tellurium is not precipitated in concentrated hydrochloric acid solution, while selenium is precipitated. Selenium prepared in this way is a red powder and tellurium a black powder. The elements exist in allotropic forms (see sul- phur, 273). When selenium is melted and allowed to solidify it is converted into an amorphous solid, which melts at 217, resembles lead in color, and has some of the physical properties of metals. It conducts electricity to a slight extent, the conductivity being markedly influenced by exposure to light. This fact has been ingeniously utilized in the invention of a photometer, which has been used to measure the relative intensities of the light given off by different stars. Selenium is used to some extent in coloring glass and enamels; it gives to them a characteristic pink or red color. Tellurium melts at 452 and on solidification is obtained as a crystalline substance, which resembles silver in color. Its metallic properties are more marked than are those of selenium. It has no uses at present. 446. Chemical Properties. Selenium and tellurium burn in the air and form dioxides, SeO 2 and TeO2. They also unite with chlorine and form compounds represented by the following for- mulas: Se 2 Cl 2 , SeCU, TeCl 2 , TeCU. Both elements react when heated with metals, and selenides and tellurides are formed. 447. Hydrides. When ferrous selenide is treated with con- centrated hydrochloric acid, hydrogen selenide is formed: FeSe + 2HC1 2 = FeCl 2 + H 2 Se The compound is a poisonous gas, which has a disagreeable, characteristic odor; it dissolves in water and shows, in general, a chemical behavior similar to that of hydrogen sulphide. Hydrogen telluride is made in a similar way and shows similar properties. Aqueous solutions of the hydrides of sulphur, selenium, and tellu- rium are decomposed slowly by the oxygen in the air, as the result of which water and the free elements are formed ; the rate at which this decomposition takes place increases with increasing atomic weight of the element. As we pass from sulphur to tellurium metallic properties become more evident and, as a result, the hydrides become less stable. The selenides and tellurides of the 394 INORGANIC CHEMISTRY FOR COLLEGES metals, with the exception of those of the alkali metals, are insol- uble in water. 448. Oxides and Oxygen Acids. The dioxides of selenium and tellurium are formed when the elements are burned in the air or oxidized with nitric acid. Selenium dioxide is a white solid, which sublimes readily and forms long needles; it is soluble in water and reacts with it to some extent to form selenous acid: SeO 2 + H 2 <= H 2 SeO 3 The salts of selenous acid, the selenites, resemble closely the sul- phites. Selenous acid is reduced by sulphurous acid to selenium: H 2 SeO 3 + 2H 2 SO 3 = Se + 2H 2 SO 4 + H 2 O This reaction is commonly used in testing for selenium because all of its compounds can be readily converted into selenous acid, and the formation of the red precipitate is characteristic. It can also be used for the quantitative determination of the element because selenium can be heated below 100 without change. Tellurium dioxide is insoluble in water, but dissolves in solu- tions of alkalies and forms salts of tellurous acid, H 2 TeO 3 , which is precipitated when soluble tellurites are treated with an acid. Tellurium dioxide shows weakly basic properties; it dissolves in concentrated hydrochloric acid and forms a chloride, TeCU, which, however, is decomposed when water is added to the solu- tion: TeCU + 3H 2 O <= H 2 Te0 3 + 4HC1 It will be recalled that the chlorides of the strongly metallic ele- ments are not decomposed by water, and that those of the acid- forming elements are decomposed, forming hydrochloric acid and the acid derived from the acid-forming elements. It is seen from the above that tellurium chloride stands between these two extremes; it is more or less decomposed, depending upon the con- centration of hydrochloric acid present. Tellurium dioxide dissolves in concentrated sulphuric acid, and the sulphate formed has the composition represented by the formula [TeO 2 ] 2 ,SO 3 . The normal sulphate of tellurium would have the formula TeO 2 ,2S0 3 ; it is seen, therefore, that the com- pound formed contains a large excess of the oxide of the basic SELENIUM AND TELLURIUM 395 element over that present in the normal salt. The compound formed is a basic sulphate (241). In general, weakly basic ele- ments form basic salts. The oxide forms a basic nitrate when dissolved in concentrated nitric acid; its formula is generally written Te2Os(OH)NO3. It can be considered as made up as represented by the following formula: [TeO2]4,H2O,N205. Tellurium dioxide and tellurous acid react with bases and form salts in which tellurium plays the part of the acid-forming element; they also react with acids and form salts in which tellurium plays the part of a base-forming element. Other elements form compounds which behave in this way and to emphasize this prop- erty such compounds are said to be amphoteric, the adjective being derived from the Greek word meaning both. Selenic acid, H 2 SeO4, can be prepared by oxidizing selenous acid with chlorine-water: H 2 SeO 3 + C1 2 + H 2 O + H 2 SeO 4 + 2HC1 The reaction is a reversible one, for in the presence of strong hydro- chloric acid selenic acid is reduced to selenous acid. Selenic acid is a more active oxidizing agent than sulphuric acid, which, it will be recalled, will not oxidize hydrochloric acid; it will oxidize gold, which is not attacked by nitric acid. Telluric acid is prepared by treating tellurous acid with chromic acid, which is a very active oxidizing agent. The formula of the acid is HeTeOe, and is written in this way and not as H2TeO4,2H 2 O in order to bring out the fact that the water is not combined as it is in a salt containing water of crystallization (water of hydration). The acid is a white crystalline solid, which when heated loses water and is converted into tellurium trioxide, TeOa, or, at higher temperatures, into the dioxide, TeO 2 . Telluric acid is readily reduced when heated with a solution of hydrochloric acid, and the tellurous acid formed is converted into tellurium tetrachloride. A careful comparison of the behavior of the compounds of selenium and tellurium with that of the analogous compounds of sulphur, will bring out clearly the relationships in physical and chemical properties which occur in a well-characterized family of elements. CHAPTER XXVIII PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 449. The elements to be considered in this chapter are mem- bers of the fifth group in the periodic classification. They show in their physical properties and chemical behavior a wide varia- tion, which, however, is systematic and progressive with changing atomic weight. There is no striking similarity between nitrogen, the first ele- ment in the group, and phosphorus, either in the elementary con- dition or in their compounds. Nitrogen is inert except at high temperatures; when it unites with oxygen the reaction is endo- thermic, and the compounds formed decompose more or less readily with the evolution of energy; nitric acid, as a result, is an active oxidizing agent. Phosphorus, on the other hand, is one of the most active elements at ordinary temperatures; it takes fire spontaneously in the air, and when it burns a large amount of heat is given off; phosphoric acid and its anhydride are very stable substances and do not act as oxidizing agents. It has already been pointed out that in most cases the first member in a group in the periodic classification differs markedly in chemical behavior from the other members of the group, and the fact was illustrated in the case of fluorine. The difference between nitrogen and phosphorus is much greater than that between fluorine and chlorine. In the case of the halogen family there is a progressive change in activity of the elements as measured by the amount of energy set free when they unite with another element. A similar change in activity is exhibited by phosphorus, arsenic, antimony, and bismuth, and they, accordingly, constitute a chem- ical family. The combining power, however, of nitrogen and phosphorus as measured by valence, is the same, since both elements have the valencies normal to all members of the fifth group, namely +5 396 PHOSPHORUS, ARSENIC, ANTIMONY AND BISMUTH 397 and 3; as a result, the compounds of the two elements have similar formulas such as NH 3 , PH 3 , N 2 O 3 , P 2 O 3 , HNO 3 , HPO 3 , etc. But a similarity in the chemical formulas of two compounds is not necessarily associated with similarity in chemical behavior. The energy relationships are the determining factor in the latter, and, as has been said, nitrogen and phosphorus differ markedly in this respect. For these reasons phosphorus is considered the first member of the chemical family in the fifth group in the periodic classification of the elements. PHOSPHOKUS 450. Occurrence. Phosphorus occurs widely distributed over the earth's surface in the form of phosphates in the soil, and is present in the so-called phospho-proteins, which are present in all living things. The bones and teeth of animals contain approxi- mately 30 per cent of calcium phosphate. The element occurs in several minerals, of which calcium phosphate (phosphorite), Ca 3 (PO4)2, and apatite, CaF 2 ,3Ca 3 (P04)2, are examples. Large beds of phosphate rock are found in the United States in Florida, Georgia, Tennessee, North and South Carolina, Utah, Montana, and Wyoming; they also occur in Ontario, Tunis, and Algeria. 451. Discovery and Preparation. Phosphorus was discovered in 1669 by Brand, an alchemist, who lived in Hamburg, Ger- many. In his search for the philosopher's stone he distilled at a high temperature the residue obtained by evaporating urine to dryness. The properties of the substance were such that it aroused a great deal of interest, and it was exhibited in the courts of Europe. When phosphorus is spread over any surface it glows in the dark with a dull light, which constantly shifts in intensity from place to place. It was this property of the element that led to its name, which was derived from the Greek words signifying light- bearer. The method of preparing phosphorus from calcium phosphate obtained from bones was first published by Scheele in 1771. In this method, which was used for many years to prepare the element on the commercial scale, calcium phosphate is heated with sul- phuric acid and a small quantity of water; the calcium sulphate 398 INORGANIC CHEMISTRY FOR COLLEGES formed is separated by filtration, and the residue, which con- tains phosphoric acid, is evaporated and distilled at a very high temperature with carbon in an earthenware retort. The phos- phorus, which passes over as a gas, is collected under water and is obtained as a yellow solid. In the industrial method of preparation now used the reac- tion is carried out in one operation in an electric furnace, as a very high temperature (white heat) is required. The calcium phosphate is decomposed by silicon dioxide (sand) which replaces the sulphuric acid used in the older method; and the phosphorus pentoxide liberated as a result is reduced to phosphorus by the carbon present. The reactions involved are indicated by the fol- lowing equations in which calcium phosphate, Cas(PO4)2, and calcium silicate, CaSiOs, are represented as made up in each case of basic and acidic oxides: (CaO) 3 ,P 2 O 5 + 3SiO 2 = 3(CaO,SiO 2 ) + P 2 O 5 P 2 O 5 + 5C = 2P + 5CO The ^apor from the furnace is cooled and led into water where the phosphorus condenses to a solid. To obtain it in a convenient form it is melted under water and cast in molds into sticks. 452. Physical Properties. Phosphorus exists in three allo- tropic modifications, which differ markedly in physical properties. White phosphorus obtained when the vapor of the element is condensed, is a horny substance which can be cut with a knife; it has the density 1.83, melts at 44, boils at 287, and is very sol- uble in carbon disulphide, and less so in ether, chloroform, and other organic liquids. Commercial white phosphorus has a light yellow color, which is due to traces of impurities. Up to about 500 the vapor of phosphorus has a density which leads to the conclusion that its molecule contains four atoms. At 1700 the molecules are partly diatomic, P2; the lowering of the freezing-point of solutions of phosphorus indicates that in this con- dition the molecule is tetra-atomic, P4. Red phosphorus is a red powder, which consists of microscopic crystals. It is prepared by heating the white variety at about 250. It is insoluble in all liquids, and does not melt when heated, but changes into vapor which is identical with that obtained from PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 399 white phosphorus. The density of red phosphorus varies between 2.05 and 2.34, and is determined by the conditions which exist when it is made. Black phosphorus has been obtained by heating the white variety at 200 under a pressure of 1200 kilograms per square centimeter, which is approximately 116 atmospheres. It is a black, lustrous substance, which has the density 2.69, and con- ducts electricity. 453. Chemical Properties. White phosphorus reacts so readily with the oxygen of the air that it is kept under water. Its kindling-point in the air is about 35, and as it oxidizes freely at lower temperatures, it is apt to take fire spontaneously. Phosphorus burns in chlorine and forms the pentachloride, PCls, and by regulating the conditions the trichloride, PCls, can be formed. It unites vigorously also with the other halogens, and combines with most elements when heated with them; with metals, phosphides are formed, of which calcium phosphide, CasP2, is an example. If white phosphorus ignites while in contact with the flesh it produces a severe burn, which heals very slowly; it should not be touched, therefore, and should be handled with great caution with the aid of pincers. The element is an active poison, about 0.15 gram being a fatal dose for man. It is used on account of this property as an ingredient of rat poisons, which contain usually, in addition to the element, a fat and a diluent such as flour. White phosphorus was formerly much employed in making matches, but as the workmen in the factories often developed a terrible disease which attacked the teeth and jaw-bone, a prohibitive tax was put on the matches manufactured in this way. Red phosphorus, which is not poisonous, is now used for this purpose. When phosphorus burns, the pentoxide formed is produced as a dense white cloud, the obscuring power of which is greater than that of a cloud produced in any other way. For this reason shells containing phosphorus were used in the recent war in setting up a cloud barrage, and in determining the accuracy of gun-fire. The puff of white smoke produced when a shell containing phos- phorus exploded could be seen at a long distance, and the place where it fell could be accurately noted. White phosphorus was also used in tracer bullets to ignite the hydrogen in war-balloons. 400 INORGANIC CHEMISTRY FOR COLLEGES When red phosphorus is formed by heating the white variety of the element, the change is accompanied by the evolution of a large amount of heat; the red form of the element is much less active, as we might expect, than the white. White phosphorus slowly changes in the light to the red form; it is for this reason that sticks of the element in the laboratory become colored on the outside, the shade being dependent on the intensity of the light and the time the element has been exposed. The kindling temperature of red phosphorus in air is about 240. Red phosphorus, which is not poisonous, enters into reactions with other elements to form the same compounds produced from the white variety. Most of the phosphorus made is used in the manu- ufacture of matches. (See problem 15, page 420.) 454. Phosphorescence. The slow oxidation of phosphorus in the air is accompanied by the giving off of a feeble light; the element is said to phosphoresce. The chemical change which takes place involves, in all probability, the formation at first of phorphorus trioxide, and the subsequent oxidation of the latter to the pentoxide. This appears to be a reasonable explanation in view of the fact that the vapor of phosphorus trioxide phos- phoresces in the air. The phenomenon occurs during the slow oxidation of phosphorus at room temperature, only when the pressure of the oxygen in contact with the element is 200 mm. or less. 465. Other substances phosphoresce and certain living things emit a light which resembles that given off by phosphorus; among the latter are glow-worms, fire-flies, and certain varieties of fish. The phosphorescence frequently observed in the wake of a steamer at sea is produced as the result of the agitation of small microscopic animals. The phenomenon of phos- phorescence is an important one about which little is known. It is produced, no doubt, as the result of the direct change of chemical energy into light. When substances burn with the evolution of light, the latter is produced as the result of the fact that the heat liberated in the chemical reactions sets the molecules in such rapid vibration that they emit light. The wave- length of the light emitted where a substance is heated is determined by the temperature attained and in most cases is the same whatever the nature of the substance. The radient energy, a part of which is light, given off from a glowing body is proportional approximately to the fourth power of the temperature (Stephan's law). There are certain substances the behavior of which is not in accord with this law. This is notably the case with phosphorus, which emits light at ordinary temperatures, and with magnesium. When the latter burns an PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 401 intense white light is given off which is utilized in photography, magnesium being one of the ingredients of flash-powders. The temperature produced when the metal burns is over 3000 below that at which a body would give off light of this character if it were the result of heat energy alone. In such cases as these a part of the chemical energy is directly changed into light energy. This kind of change is utilized when gas is burned in connection with a Welsbach gas-mantle. The materials of which the mantle is made apparently facilitate the change of chemical energy into visible light, for the temperature of the mantle is far below that necessary to produce the quality of light obtained. The extent of this change is very small, however, compared with that of the change of chemical energy into heat. 456. Oxides of Phosphorus. When phosphorus is burned in air or oxygen, phosphorus pentoxide, P2Os, is formed. The com- pound is a white powder, which can be melted. It reacts with water violently and is converted into metaphosphoric acid: P 2 O 5 + H 2 O = 2HPO 3 On account of this property phosphorus pentoxide is used as a drying agent for liquids and gases with which it does not react. It is the most efficient agent known for this purpose. The oxide reacts with many compounds containing hydrogen and oxygen, and extracts from them these elements in the proportion required to form water; it is thus a dehydrating agent. Its behavior in this way with nitric acid and sulphuric acid has been noted; in each case the anhydride of the acid is formed, along with meta- phosphoric acid. The pentoxide will remove hydrogen and oxygen from paper, which is made of cellulose, (CeHioOs)*, and the carbon left behind will be evident as the result of the black color produced when the oxide came in contact with the paper. When phosphorus pentoxide is heated with water, the meta- phosphoric acid first formed is converted into phosphoric acid, HsPC^; for this reason the pentoxide is called phosphoric anhy- dride. 457. The trioxide of phosphorus, P2Os, is formed when the element is oxidized in a limited supply of air. It is a white solid which phosphoresces in air, and oxidizes to the pentoxide. It melts at 22.5 and boils at 173. Phosphorus trioxide reacts with water at ordinary temperatures to form phosphorous acid, HsPOs. 458. The Phosphoric Acids. Three well-characterized acids are derived from the single anhydride, P205. Since they are 402 INORGANIC CHEMISTRY FOR COLLEGES closely related and each contains phosphorus with the valence 5, they are called phosphoric acids, and a prefix is- added to the name to designate each one. Metaphosphoric acid has the formula HPO 3 (P205,H 2 O), pyrophosphoric acid H 4 P 2 O 7 (P205,2H2O), and orthophosphoric acid H3PO 4 (P2O5,3H2O). The relationship exist- ing between the acids is clearly brought out by indicating their composition in the way shown in the parentheses. When phos- phorus pentoxide is dissolved in water, metaphosphoric acid is first formed. On standing or on heating the solution, further hydration takes place and orthophosphoric acid is produced. When the latter is obtained by evaporating the water and is heated for some time at 255 it is converted into pyrophosphoric acid. Meta- phosphoric acid is called glacial phosphoric acid in commerce, and is usually obtained in the form of transparent sticks. 459. Orthophosphoric Acid. The calcium salt of this acid occurs in nature and is used as the commercial source of the acid, which is prepared by treating the mineral with sulphuric acid. The acid may be prepared by oxidizing red phosphorus with strong nitric acid. It is a white solid which melts at 42.3. Orthophosphoric acid, HsPO^ contains 3 hydrogen atoms which can be replaced by metallic atoms; it is, consequently, tribasic. When dissolved in water it ionizes as indicated by the following formulas : H 3 P0 4 + H + + H 2 P0 4 ~ <=* H + + HP0 4 ' 5 H + + PO 4 ~ In 0.1N solution the first step in the ionization takes place to the extent of about 27 per cent at 18; the bivalent and trivalent ions are formed to a much smaller extent. Orthophosphoric acid is, accordingly, a much weaker acid than sulphuric acid, which is ionized to the extent of 61 per cent at the same concentration. 460. Orthophosphates. Orthophosphoric acid, like other tri- basic acids, forms three classes of salts which are produced as the result of the replacement of one-third, two-thirds, and all the hydrogen in the acid. The formulas of the sodium salts are as follows: NaH 2 PO 4 , Na 2 HPO 4 , and Na 3 PO 4 ; the first is called monosodium phosphate or primary sodium phosphate, the second disodium phosphate or secondary sodium phosphate, and the third trisodium phosphate or tertiary sodium phosphate. PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 403 461. The primary salt shows a slight acid reaction; the ions chiefly formed from the salt are Na + and H 2 PO 4 ~, but the latter ionizes slightly to form H + and HPO 4 ~~, and, as a result, the solu- tion shows an acid reaction. The secondary salt shows a very weak alkaline reaction as the result of hydrolysis. In this case the ions formed in largest amount are 2Na + and HPO 4 ~~, but the latte'r reacts to some extent with the hydrogen ions formed from water, H 2 <= H + -f OH~, and in uniting with them to form H 2 PO 4 ~(HPO 4 ~ + H + <=H 2 PO 4 ~) leaves an ex- cess of OH~ ions in the solution, which, as a result, shows an alka- line reaction. The tertiary salt, which can be made by evaporating to dryness a solution of the secondary salt and sodium hydrox- ide, is converted by water into the secondary salt and sodium hydroxide : Na 3 PO 4 + H 2 O <=> Na 2 HPO 4 + NaOH The equation written with ionic formulas to emphasize the hydroly- sis is as follows : 3Na + + PO 4 - ' + H + OH~ <=3Na + HPO 4 ~ + OH~ 462. The action of heat on primary and secondary phosphates is represented by the following equations : NaH 2 PO 4 = H 2 O + NaPO 3 2Na 2 HPO 4 = H 2 O + Na 4 P 2 O 7 In both cases the elements of water are lost ; primary orthophos- phates are thus converted into metaphosphates, and the secondary salts into pyrophosphates. The reactions take place very slowly in the reverse direction in aqueous solutions of metaphosphates and pyrophosphates. When sodium-ammonium phosphate (microcosmic salt) is heated, both water and ammonia are lost: NaNH 4 HPO 4 = NH 3 + H 2 O + NaPO 3 Magnesium-ammonium phosphate loses ammonia under the same conditions: 2MgNH 4 PO 4 = 2NH 3 + H 2 + Mg 2 P 2 7 404 INORGANIC CHEMISTRY FOR COLLEGES The first-named salt is used in qualitative analysis (see borax- beads), and the latter in quantitative analysis. Magnesium- ammonium phosphate, which is insoluble in water, is the salt commonly used to separate magnesium from other metals in qualitative and quantitative analysis. When it is used for the latter purpose it is heated and weighed as the pyrophosphate. The normal phosphates, except those of the alkali metals, are insoluble in water, but dissolve in strong acids; the acid phosphates of the alkaline earths (calcium, strontium, etc.) are also soluble in water. 463. Tests for Orthophosphates. Orthophosphoric acid and its soluble salts are converted in solution by silver nitrate into silver orthophosphate, Ag3PO4, which is formed as a yellow precipitate. This test serves to distinguish the acid and its salts from the other phosphoric acids or their salts, which give white precipitates under the same conditions; it is not, however, characteristic of orthophosphates. Another test, frequently used, consists in adding to the solution of the acid or salt a solution of magnesium chloride and ammonia; magnesium-ammonium phosphate is formed as a crystalline precipitate. A third test commonly applied in qualitative and quantitative analysis is based on the formation of a copious yellow precipitate when a solution of a phosphate is treated with a solution of ammo- nium molybdate, (NH^MoQ*, containing nitric acid. The precipitate has the complex composition represented by the formula (NH4)3PO4,12MoO 3 ,6H 2 O, and is called ammonium phos- phomolybdate. Orthoarsenic acid and its salts, which resemble closely the analogous compounds of phosphorus, give a precipitate similar in appearance and composition to that derived from ortho- phosphoric acid; for this reason positive tests for orthophosphates should be followed by a test for arsenic. Silver or thoar senate is a brown precipitate. 464. Phosphorous Acid, H 3 PO 3 . The formation of this acid from phosphorus trioxide has already been noted. It is best pre- pared by the action of phosphorus trichloride on water: PC1 3 + 3H 2 = H 3 P0 3 + 3HC1 Phosphorous acid is a comparatively weak acid and forms salts as the result of the replacement of one or two hydrogen atoms only. PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 405 It is an active reducing agent and precipitates the noble metals from solutions of their salts. When heated it decomposes, and metaphosphoric acid and phosphine, PH 3 , are formed, the change in valence being from +3 in phosphorous acid to +5 in meta- phosphoric acid and 3 in phosphine. 465. Hypophosphorous Acid, H 3 PO 2 . Salts of hypophosphor- ous acid are formed when phosphorus is dissolved in a solution of a strong base. It will be recalled that when chlorine reacts with sodium hydroxide, sodium hypochlorite and- sodium chloride are formed, and the reaction was explained on the assumption that the halogen reacted with water and formed hypochlorous acid and hydrochloric acid, which were subsequently neutralized by the base present (432). It is probable that a similar reaction takes place in the case of phosphorus, as a result of which hypophos- phorous acid and phosphine are produced: 4P + 6H 2 O = PH 3 + 3H 3 PO 2 Phosphine is not an acid and is not affected, therefore, by the base present and escapes as a gas; the acid produced, which is monobasic, is converted into a salt. The equation for the reaction when white phosphorus is dissolved in a warm solution of sodium hydroxide is, therefore, as follows: 4P + 3H 2 O + 3NaOH = PH 3 + 3NaH 2 PO 2 In preparing the free acid, barium hydroxide is used instead of sodium hydroxide, and the barium salt formed is converted into the acid and barium sulphate by adding sulphuric acid to the solution; on removing the precipitate by filtration a solution of hypophos- phorous acid is obtained. The acid is a white crystalline solid; it is a very active reducing agent. 466. Phosphine. -A reaction by which phosphine can be pre- pared from white phosphorus and a solution of a base has just been discussed (465) in some detail. Certain side-reactions occur which lead to the formation of hydrogen and a second hydride of phosphorus of the formula P 2 H4, which is a liquid that boils at 57. Since the latter is spontaneously inflammable, the gas generated takes fire when it comes into contact with the air, unless it has been freed from the liquid hydride by passing it through alcohol or some other solvent. 406 INORGANIC CHEMISTRY FOR COLLEGES Phosphine can be prepared by decomposing with water the phosphides of the more active metals; calcium phosphide can be used conveniently for this purpose : Ca 3 P 2 + 6H 2 O = 3Ca(OH) 2 + 2PH 3 Phosphine is a colorless ; inflammable gas. The products of its combustion are water and metaphosphoric acid, which is produced as the result of the union of the phosphorus pentoxide formed with a part of the water produced. Phosphine resembles ammonia in composition and like ammonia unites with the halogen hydrides to form compounds analogous in composition to ammonium salts; the most stable of these is phosphonium iodide, PELiI. The phosphonium compounds are decomposed by water into phosphine and the free acid. Phosphine is insoluble in water and does not react with it to form a compound analogous to ammonium hy- droxide. Phosphine resembles, in some respects, hydrogen sulphide; when passed into solutions of the salts of some of the metals it causes the precipitation of insoluble phosphides; the case of copper is an example : 3CuSO 4 + 2PH 3 = Cu 3 P 2 + 3H 2 SO 4 467. Halides of Phosphorus. Phosphorus trichloride, PC1 3 , is formed when chlorine is passed over white phosphorus, care being taken to avoid an excess of the halogen; under the latter circum- stances the pentachloride, PCls, is formed, since it is produced as the result of the direct action of the trichloride and chlorine. The trichloride is a heavy, colorless liquid, which boils at 76. It reacts with water to form phosphorous acid and hydrochloric acid, and in the hydrolysis resembles the halides of other acid-forming elements. Phosphorus pentachloride is a white solid, which sublimes at 163, and dissociates into the trichloride and chlorine at higher temperatures. It is hydrolyzed very rapidly by water, the product formed being determined by the amount of water present. With excess of water complete hydrolysis takes place and phosphoric acid and hydrochloric acid are formed : PC1 5 + 4H 2 O = H 3 PO 4 -f 5HC1 PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 407 If not enough water is present to bring about this reaction partial hydrolysis takes place and phosphorus oxychloride is formed: PC1 5 + H 2 O = POC1 3 + 2HC1 Phosphorus forms with fluorine and with bromine compounds in which it shows the valencies 3 and 5. The two iodides have the formulas P2l4 and Pis. All these compounds resemble the chlo- rides in chemical behavior and are prepared by the same general methods as those used to prepare the latter. 468. Sulphides of Phosphorus. Compounds represented by the following formulas can be prepared by heating phosphorus with the proper amounts of sulphur: P 4 S 3 , P4S 7 , and P 2 Ss They are all solids which are hydrolyzed by water. Phosphorus pen- tasulphide reacts with the latter and forms phosphoric acid and hydrogen sulphide: P 2 S 5 + 8H 2 O = 2H 3 PO 4 + 5H 2 S The sulphide of the formula P 4 S 3 is used in making matches. ARSENIC 469. The chemistry of arsenic and its compounds resembles very closely that of phosphorus and its compounds. Since the latter has just been discussed at length it will be possible to limit largely the consideration of arsenic to an account of any unique properties exhibited by it or its compounds, and to the uses to which they are put. Arsenic is a much less active element than phosphorus and is less strongly electro-negative in character. Its compounds with hydrogen and with oxygen are less stable than the corresponding phosphorus derivatives. 470. History and Occurrence. Two sulphides of arsenic occur as minerals which are highly colored and, therefore, early attracted attention; orpiment, As 2 S 3 , is bright yellow, and realgar, As2S 2 , is orange-red. When heated in the air both compounds are con- verted into arsenic trioxide, As 2 O 3 , which occurs also as a mineral called arsenite. The properties of these substances were known to the early Greeks, who gave to orpiment a name from which the word arsenic is derived. The element in the free condition was 408 INORGANIC CHEMISTRY FOR COLLEGES known to the alchemists in the thirteenth century, a fact which resulted, no doubt, from the ease with which arsenic oxide is reduced when heated with charcoal. The poisonous properties of the oxide have been known and used since the earliest times. Arsenic occurs free in nature, in combination with metals as arsenides such as smaltite, CoAs2, and in association with sulphur in metallic sulphides. Arsenical pyrite, mispickel, FeAsS, is an important mineral, which occurs along with pyrite, FeS 2 ; zinc sulphide, sphalerite, ZnS, also contains compounds in which the sulphur is in part replaced by arsenic; cobaltite has the formula CoAsS. When pyrite is burned to make sulphur dioxide in the manufacture of sulphuric acid, the arsenic in it is converted into arsenic trioxide, and since the latter is a solid it condenses in this form in the dust boxes and flues. In the smelting of zinc ores containing arsenic, the latter is obtained as the trioxide when the ores are heated in the air (roasted) preparatory to reduction. The oxide obtained from these sources is heated with carbon and the element liberated is distilled off and condensed. Arsenic is also obtained by heating arsenical pyrites; arsenic distills off and fer- rous sulphide is left in the retort. 471. Physical Properties. Arsenic has the appearance of a metal; it is steel-gray in color, is more or less brittle, and when fractured, shiny crystalline surfaces can be seen. It has the den- sity 5.7, and is a poor conductor of electricity. Arsenic sublimes at relatively low temperatures, and at 600 the pressure of its vapor is equal to 1 atmosphere. The vapor has a light-yellow color and its density indicates that its formula is As4; when cooled quickly it condenses to a yellow modification of the element, which resembles closely white phosphorus in properties, being phos- phorescent and soluble in carbon disulphide. 472. Chemical Properties. When arsenic burns in air or oxy- gen the trioxide, As2Os, is formed. It unites with the halogens and sulphur readily and with the metals forms arsenides. Active oxidizing agents convert it into arsenic acid, HaAsCU. The ele- ment does not displace hydrogen from acids. 473. Oxides of Arsenic. Arsenic trioxide, As2C>3, which is formed when arsenic burns or when arsenides are heated in the air, has been known for a long time and is commonly called " arsenic " or white arsenic. It is purified by sublimation and is obtained PHOSPHORUS, ARSENIC, ANTIMONY. AND BISMUTH 409 as a white crystalline powder. If the vapor is cooled rapidly a transparent form resembling glass is produced, which passes slowly into the crystalline variety. The oxide dissolves to a slight degree in hot water and reacts with the latter to form arsenious acid, HaAsOa. It is more soluble in strong hydrochloric acid with which it forms arsenic trichloride; with concentrated sulphuric acid, a basic sulphate is formed, which is hydrolyzed by water. Arsenic trioxide is an active poison; the fatal dose is from 0.06 to 0.18 gram (1 to 3 grains). Many people who live in high alti- tudes are said to eat white arsenic as it is a help in respiration, which is more or less difficult in places where the concentration of the oxygen in the air is low. A solution of the oxide in sodium carbonate (Fowler's Solution) is administered as a stimulant in certain nervous affections. Arsenic trioxide is used in making high-quality colorless glass; it serves as an oxidizing agent, the arsenic produced being vola- tilized. It is used to some extent as a mordant in calico-printing, in making pigments, as an ingredient of poisons for rats and flies and other insects, and as a preservative for untanned hides. Arsenic pentoxide is obtained as a white crystalline substance by cautiously heating arsenic acid. It dissolves in water and reacts with it to form arsenic acid. 474. Acids of Arsenic. Salts containing arsenic are known which resemble in composition and physical and chemical proper- ties the salts of ortho-, pyro-, and meta-phosphoric acids and phosphorous acid. Orthoarsenic acid, (H3AsO4)2,H2O, is the only acid of arsenic which has been isolated. Silver orthoarsenate, which is obtained as a brown precipitate, and magnesium-ammo- nium arsenate are characteristic salts that resemble the corre- sponding salts of orthophosphoric acid. Arsenious acid is a very weak acid and its soluble salts, which are formed by dissolving arsenic trioxide in solutions of the caustic alkalies, are highly hydrolyzed ; the arsenites of the heavy metals prepared from these by double decomposition possess, in most cases, complex for- mulas. 475. Antidotes for Poisons. When a poison has been taken into the body through the mouth, the usual procedure is to admin- ister as soon as possible an emetic, which causes vomiting. An antidote is next taken, the object of which is to convert the poison 410 INORGANIC CHEMISTRY FOR COLLEGES into an insoluble substance and thus reduce as much as possible the absorption of the poison by the tissues of the body. In the case of white arsenic or arsenites, freshly precipitated ferric hydrox- ide or magnesium hydroxide is commonly used on account of the fact that ferric arsenite and magnesium arsenite are insoluble salts. The frequent use of white of egg as an antidote is based on the fact that the salts of the heavy metals which are poisonous form, in most cases, insoluble precipitates with the proteins present in the egg. 476. Arsine. The arsenides of the more active metals react with hydrochloric acid and form chlorides and arsine, AsHa. Sodium arsenide and magnesium arsenide are hydrolyzed by water and, thus, resemble the analogous nitrides and phosphides. Arsine can be prepared in these ways. It is a gas with a dis- agreeable garlic-like odor, and is very poisonous. It liquefies at -40. Arsine is formed when compounds containing arsenic are reduced by nascent hydrogen, and it is this method which is used in testing for arsenic, especially when it is present in small quantities in the substance under examination. Arsine is readily decomposed by heat into the elements of which it is composed, and when the decomposition is effected in a glass tube arsenic is deposited in the form of a brown coating, or as a mirror if enough of the element is present. This method of detecting arsenic is known as Marsh's test, and is the one commonly used in col- lecting evidence to be presented in court in cases of suspected poisoning by arsenic. In carrying out the test, pure zinc and pure dilute sulphuric acid are placed in a hydrogen generator, and the gas formed passed first through a tube containing dehydrated cal- cium chloride, which is used as a drying agent, and then through a tube of hard glass which is constricted at one place. When hydro- gen is freely evolved the hard glass tube is heated with a Bunsen flame at a point between the generator and the constricted part of the tube. If, after a few minutes, no deposit is formed in the hot tube, the solution to be tested for arsenic is poured through the thistle tube into the generator. The preliminary heating is necessary, because most samples of zinc and of sulphuric acid con- tain traces of arsenic, which are present as the result of the fact that both of these substances are made from raw materials that PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 411 are apt to contain arsenic. If the compound introduced into the generator contains arsenic, it is reduced by the nascent hydrogen; the arsine formed is decomposed in the hot part of the tube and arsenic is deposited in the constriction of the latter. In using this method to determine the arsenic quantitatively, a series of tubes is prepared in which arsenic is deposited from known amounts of a solution of sodium arsenite. By comparing the size and appearance of the deposit on these standards with the tube used in the case of the substance being tested, it is possible to determine the quantity of arsenic present. When compounds containing antimony are heated with nas- cent hydrogen a volatile hydride, easily decomposed by heat, is formed. It is possible, however, to distinguish the deposit obtained in this case from that produced by arsenic (487). Arsine burns when ignited, and is converted into water and arsenic trioxide, which forms a white cloud. If a piece of porce- lain is put into an arsine flame, a black deposit of arsenic is formed on the cold surface, the phenomenon observed being similar to that which occurs when a cold object is brought into contact with an illuminating gas flame (195). Arsine is insoluble in water, does not react with aqueous solu- tions of acids to form salts as ammonia does, and does not exhibit the behavior shown by phosphine in combining with the halogen hydrides. 477. Halides of Arsenic. The compounds of arsenic and the halogens have the general formula AsXs. They are hydrolyzed by water, but owing to the slight basic properties of arsenious acid the reactions are reversible. Arsenic trichloride boils at 130, and when a solution of it in hydrochloric acid is distilled, it passes over with the acid and the water; it is thus possible to separate arsenic in this way from many other elements. 478. Sulphides of Arsenic. The occurrence of two sulphides of arsenic as minerals has already been noted. Artificial orpiment, As2S3, is made by subliming a mixture of arsenic trioxide and sul- phur and is used as a yellow pigment. Realgar, As2S2, is made by fusing together the oxide and sulphur, or by distilling arsenical ores with sulphur. It is used to some extent as a red pigment, but its chief use is in making mixtures to produce colored lights when burned, and for removing the hair from hides preparatory to tanning. 412 INORGANIC CHEMISTRY FOR COLLEGES Arsenic trisulphide is formed as a yellow precipitate when hydrogen sulphide is passed into a solution of arsenic trichloride. It is separated in this way in the procedure used in qualitative analysis. The sulphide is practically insoluble in hot concentrated hydrochloric acid, a fact which is used in separating it from anti- mony trisulphide, which dissolves in this reagent. Arsenic trisulphide dissolves in ammonium sulphide as the result of the formation by direct addition of a compound, which is soluble in water: As 2 S 3 + 3(NH 4 ) 2 S = As 2 S 3 ,3(NH 4 )2S The formula of the compound is generally written as (NH 4 ) 3 AsS 3 , and it is called ammonium thioarsenite; it can be considered as the ammonium salt of the acid H 3 AsS 3 formed as the result of the replacement of the oxygen in arsenious acid, H 3 AsO 3 , by sul- phur. Many salts of such thioacids are known. When an acid is added to the solution of ammonium thioarsenite the salt is decomposed: (NH 4 ) 3 AsS 3 + 3HC1 = 3NH 4 C1 + H 3 AsS 3 2H 3 AsS 3 = As 2 S 3 + 3H 2 S Arsenic trisulphide is precipitated and hydrogen sulphide is set free. 479. When arsenic trisulphide is treated with a solution of yellow ammonium sulphide, which contains the polysulphide (284), some of the sulphur of the latter unites with the trisulphide to form the pentasulphide of arsenic, which, in turn, reacts with ammonium sulphide to form the compound As 2 Sr,,3(NH 4 ) 2 S. The formula of this compound is written (NH 4 ) 3 AsS 4 and it is called ammonium thioarsenate; it bears the same relation to ammonium arsenate, (NH 4 ) 3 AsO 4 , that ammonium thioarsenite, (NH 4 ) 3 AsS 3 , bears to ammonium arsenite, (NH 4 ) 3 AsO 3 . When a solution of the thioarsenate is treated with an acid, arsenic pentasulphide is precipitated : (NH 4 ) 3 AsS 4 + 3HC1 = 3NH 4 C1 + H 3 AsS 4 2H 3 AsS 4 = As 2 S 5 + 3H 2 S These reactions are used in qualitative analysis. The salts are often called sulpharsenites and sulpharsenates. PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 413 When hydrogen sulphide is passed into a solution of arsenic acid in concentrated hydrochloric acid, arsenic pentasulphide is precipitated. ANTIMONY 480. A comparison of the chemistry of antimony with that of phosphorus and arsenic brings out clearly the progressive change in properties within a chemical family, with increase in atomic weight. Phosphorus acts as an acid-forming element only. Arsenious hydroxide is both a base and an acid. When the hydroxide of an element shows both acidic and basic properties the element is said to be amphoteric. We shall see that antimonous acid is a weaker acid than arsenious acid, but that it is a stronger base than the latter, since its salts cannot be completely hydrolyzed by cold water. The pentoxide of antimony acts as an acid anhy- dride only. This fact serves as an example of the generalization that an increase in the valence of an element toward oxygen is asso- ciated with an increase in acidic properties, and, as a result, a decrease in basic properties. Of two acids derived from the same element the stronger acid is the one in which the element shows the higher valence. There is a gradation in the properties of the hydrides of the elements in the phosphorus family. Since hydrogen acts as a posi- tive element, the most stable hydrides are those in which it is com- bined with strongly negative elements. As the negative character of the elements in the phosphorus family drops off, the stability of the hydrides decreases; antimony hydride is decomposed into its elements at a much lower temperature than is arsine. Similar relationships exist among other compounds of these elements, which it will be of interest to the student to discover. 481. History. Compounds of antimony have been known since Biblical times, and since they were used for a number of pur- poses they were studied in considerable detail by the alchemists and by their successors, the so-called iatrochemists, who centered their attention on the use of chemical substances as drugs. The effect on the body of many inorganic compounds and organic substances isolated from plants was studied by the iatrochemists in the Middle Ages, and the facts discovered became the basis of the branch of medicine called materia medica, which has to do 414 INORGANIC CHEMISTRY FOR COLLEGES with the materials used in the science on account of their effect on the body. Many of the substances prescribed to-day as drugs were first studied by the iatrochemists, of whom Paracelsus was the leader. Basil Valentine first described in the fifteenth century a method of preparing antimony. 482. Occurrence. Antimony occurs free in small quantities, usually in association with arsenic. Its chief ore is stibnite, Sb2Sa, which is found in the form of large glistening black prisms. Two oxides, Sb2Os and SboCU, occur as minerals. The element, like arsenic, is associated with sulphur in certain metallic sulphides, and is obtained as a by-product in the metallurgy of these metals, of which lead and copper are examples. 483. Preparation. Antimony is obtained from stibnite either by direct reduction with iron, or by first roasting the mineral to convert it into the oxide, and then reducing the latter with carbon. In the first process the ore is mixed with salt and scrap iron in a large crucible and heated to a high temperature in a furnace. The salt melts and serves as a flux (521), and the heavy ferrous sulphide sinks to the bottom of the crucible. The reaction which takes place is represented by the following equation : Sb 2 S 3 + 3Fe = 2Sb + 3FeS The antimony formed is poured off and cast into a convenient form. Since it contains several per cent of iron, it must be purified ; this is done by melting it with just enough stibnite to react with the iron present. 484. Properties. Antimony is a bluish-white, highly crystal- line substance with a marked luster. It can be readily powdered, and has a high density, 6.7; it melts at 630 and boils at 1440. Its vapor just above its boiling-point is made up of molecules containing 3 and 4 atoms of the element. The chief use of anti- mony is in the preparation of alloys such as type-metal and brit- tania metal, which are to be cast in molds. Owing to the presence of antimony these alloys expand on solidification and give sharp castings. Antimony burns in the air and forms the trioxide, and with chlorine and bromine it forms trihalides. It combines directly with sulphur, arsenic, phosphorus, and some metals. It is oxidized by concentrated nitric acid to the trioxide, which on continued heating PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 415 with the acid is changed to antimonic acid. It is converted by hot concentrated sulphuric acid into antimony sulphate, 802(804)3. 485. Oxides and Salts of Antimony. When antimony is burned in the air the chief product is the trioxide, SboOa, which is white; if an excess of oxygen is used the tetroxide, Sb2O4, is formed. The latter possesses neither acidic nor basic properties. Anti- mony pentoxide, Sb2O 5 , is obtained as a yellow powder when anti- monic acid is heated; it dissolves in bases and forms antimonates. Antimony trioxide dissolves in alkalies and antimonites are formed. It reacts with acids to form salts; it is converted, for example, by strong hydrochloric acid into antimony trichloride, SbCls, by nitric acid into antimony nitrate, Sb(NO3)3, and by sulphuric acid into the sulphate, Sb2 (804)3. In all these reactions the oxide shows basic properties and anti- mony acts as a metallic element. It is, however, weakly basic only, and, as a consequence, the salts are hydrolyzed by water. When water is added to a strong solution of the chloride, antimony oxy chloride is precipitated: SbCl 3 + H 2 O <=> SbOCl + 2HC1 The reaction is a reversible one and the concentrations at equilib- rium are determined by the amount of acid present in the solution. A large number of basic salts of antimony are known which contain the element and oxygen in combination as the group SbO. For this reason the latter has been given a special name and is called antimonyl; it may be considered as having the valence 1. Tartar-emetic, which has been used in medicine for a long time, is prepared by treating cream of tartar, acid potassium tartrate, KH^EUOe), with antimony trioxide; the acidic hydrogen is replaced by the univalent SbO group, and the salt formed, KSbO^HUOe), ^fbO, is potassium antimonyl tartrate. Tartar- emetic is extensively used as a mordant in dyeing vegetable fabrics, such as cotton and linen. 486. Acids of Antimony! When bases are added to solutions of antimony salts, a hydroxide, which is antimonous acid, Sb(OH)3, is precipitated. The acid soon loses water and passes into the trioxide. Antimonic acid is formed by oxidation of antimony with nitric acid, or as the result of the hydrolysis of antimony pentachloride, 416 INORGANIC CHEMISTRY FOR COLLEGES It is a white insoluble compound, which dissolves in solu- tions of the alkalies. The best known salts of antimonic acid are derived from the meta and pyro acid. Sodium pyroantimonate, Na2HoSb2O7, is difficultly soluble in water: this fact is noteworthy, as most sodium salts are readily soluble in water. 487. Stibine. The methods of preparation and properties of stibine, SbHs, resemble closely those of arsine. The deposit of the element obtained in the Marsh test, or when a cold object is brought into contact with a flame of burning stibine, differs in appearance and properties from that produced by arsenic. The deposit in the case of arsenic is shiny and brownish-black; it dissolves in a solution of bleaching powder, and is readily volatil- ized at the temperature of the Bunsen flame. The deposit of antimony is dull black, is insoluble in bleaching powder, and does not volatilize at as low a temperature as arsenic. Stibine boils at -17 and freezes at -88. 488. Halides of Antimony. Antimony trichloride was called by the alchemists " butter of antimony " on account of the fact that its crystals form a pasty mass in the air. It is hydrolyzed by water and yields several basic chlorides, the best characterized of which has the composition SbOCl. Antimony pentachloride is a heavy, colorless liquid which fumes in the air; it dissolves in water, from which it can be ob- tained in the form of hydrated crystals. Compounds of anti- mony with bromine, iodine, and fluorine are known. 489. Sulphides of Antimony. The trisulphide occurs in nature as stibnite, Sb2Ss, which is a black mineral. It is formed as an orange-red precipitate when hydrogen sulphide is passed into a solution of an antimony salt. The product prepared in this way is used as a pigment and for vulcanization in making the red antimony rubber of commerce. Antimony trisulphide dissolves in ammonium sulphide and in ammonium polysulphide, the reactions being analogous to those in the case of arsenic trisulphide. When the thioantimonate is treated with an acid, antimony pentasulphide, Sb2Ss, is formed as an orange-red precipitate; it is also formed when hydrogen sul- phide is passed into a solution of an antimonate in hydrochloric acid. Antimony trisulphide is soluble in concentrated hydrochloric PHORPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 417 acid, whereas arsenic trisulphide is not; the reaction serves, thus, to separate the two elements : Sb 2 S 3 + 6HC1 + 2SbCl 3 + 3H 2 S Antimony pentasulphide is also soluble in concentrated hydro- chloric acid, the trichloride, sulphur, and hydrogen sulphide being formed. BISMUTH 490. Bismuth is more metallic in character than antimony. It does not form a hydride and its trioxide is not soluble in bases and does not act as an acid anhydride. Bismuth trioxide dissolves in acids to form salts, which are more or less hydrolyzed by water. Bismuth pentoxide, Bi2Os, has very weak acidic properties. 491. Occurrence and Preparation. Bismuth is a compara- tively rare element; it occurs in the free condition and as the triox- ide, Bi2Os, (bismuth ocher), and as the trisulphide, 61283, (bis- muth glance), associated with silver, cobalt, nickel, and arsenic ores. The element has been known for centuries and was described by Basil Valentine. The metal is usually obtained by first roasting the ore in the air, as the result of which the sulphur and most of the arsenic present are converted into oxides and volatilized. The product is next heated with carbon (coal), iron, and a flux, and the metallic bismuth formed by reduction is drawn off while still in the liquid condition. To remove the arsenic and antimony mixed with the metal, it is fused with sodium carbonate and potassium nitrate; the latter oxidizes the arsenic and antimony and converts them into salts, but does not affect the bismuth. 492. Properties. Bismuth has a silvery luster with a reddish tinge; it is very brittle, melts at 269, boils at 1420, and has the density 9.82. Its molecular formula is Bi2, but at very high tem- peratures the molecule undergoes partial dissociation into atoms. It expands when passing from the liquid to the solid condition, whereas the common behavior is the reverse in the case of metals. Bismuth is used in making the so-called fusible alloys. When fused with other low-melting metals, a mixture is obtained which melts at a much lower temperature than that at which any of its constituents melts. Wood's metal, for example, melts at 60.5; it 418 INORGANIC CHEMISTRY FOR COLLEGES is made by melting together 4 parts of bismuth, 2 parts of lead, 1 part of tin, and 1 part of cadmium. It is possible by varying the composition of the mixtures to obtain products that melt at any desired temperature. The safety plugs used in boilers, elec- tric fuses, automatic sprinklers, and fire-doors are made of alloys prepared in this way. Bismuth unites with the halogens to form salts of the general formula BiXs, and is converted into salts when heated with oxi- dizing acids. 493. Oxides of Bismuth. The most important oxide is the trioxide, Bi2Os, which can be prepared by burning the element or by heating to a high temperature the hydroxide or nitrate of the metal; it is a yellow powder which dissolves in strong acids to form salts. When bismuth hydroxide, suspended in water, is treated with a solution of stannous chloride, SnCk, a black pre- cipitate is formed, to which the formula BiO is assigned. If the hydroxide is treated with chlorine-water, oxides are formed which have the composition Bi2O4 and Bi2Os. 494. Salts of Bismuth. Bismuth chloride, BiCl 3 ,H 2 O, can be prepared by dissolving the oxide in an excess of hydrochloric acid and evaporating the solution to crystallization. Bismuth nitrate, Bi(NOs)3,5H2O, is prepared in a similar way. Both salts are hydrolyzed by water; in the case of the chloride two chlorine atoms are replaced by hydroxyl groups: BiCl 3 + 2H 2 O ^ Bi(OH) 2 Cl + 2HC1 A similar reaction takes place in the case of the nitrate. The basic chloride is converted into bismuth oxychloride, BiOCl, when dried. The basic nitrate, Bi(OH)2NO3, is used in medicine, under the name subnitrate of bismuth, as an internal remedy in the case of certain stomach and intestinal troubles. It is also used in face-powders. Bismuth sulphide, Bi2Ss, is obtained as a brownish-black pre- cipitate when hydrogen sulphide is passed into a solution of a bis- muth salt; the presence of a dilute acid does not interfere with its precipitation. The sulphide is not soluble in ammonium poly- sulphide and can be separated by means of this reagent from arsenic and antimony. The insolubility of the sulphide is due to the fact that bismuth is more metallic in character than the other PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 419 members of the phosphorus family and does not exhibit the prop- erty of acting as an acid-forming element in compounds analogous to the thioarsenites and thioarsenates. 495. Tests for Bismuth. In the usual procedure followed in qualitative analysis, bismuth is precipitated as the sulphide from a hydrochloric acid solution along with the other elements which form sulphides insoluble in acids. When the sulphides are treated with warm ammonium polysulphide to remove the sul- phides of arsenic, antimony, and tin, bismuth sulphide is not affected as it is insoluble in this reagent. When the insoluble sulphides left after the treatment are heated with nitric acid, bis- muth sulphide dissolves and is converted into bismuth nitrate. Addition of ammonium hydroxide to the solution causes the pre- cipitation of white bismuth hydroxide, which is filtered off and dissolved in the smallest possible amount of concentrated hydro- chloric acid. The solution of bismuth chloride so formed is next poured into water, when bismuth oxychloride precipitates as the result of the hydrolysis of the chloride. EXERCISES 1. Write the equations for the reactions which take place between the following substances: (a) P and HNO 3 , (6) Ca 3 (PO 4 )2 and H 2 SO 4 , (c) PH 3 and O 2 , (d) Na and P, (e) Mg and P, (/) Al and P, (g) P 2 O 3 and H 2 O, (A) P 2 O 6 and HNO 3 , (i) P 2 O 6 and H 2 SO,, (j) P 2 O S and C 6 H 10 O 6 , (/c) H 4 P 2 O 7 and H 2 O, (I) POC1 3 and H 2 O, (ro) P and KC1O 3 . 2. How could you separate from each other the following: (a) P 2 O 3 and P 2 O 6 , (6) Ca 3 (PO4) 2 and Na 2 HPO 4 , (c) Sb 2 S 3 and Bi 2 S 3 , (d) Sb ? S 3 and 3. Write graphic formulas for the following showing in each case the-h or - valence of each element: (a) P 2 O 6 , (6) PH 3 , (c) H 3 PO 4 , (d) PC1 6 . Write two possible graphic formulas for (e) H 3 PO S and two for (/) H 3 PO 2 , and indicate in each the valence of phosphorus. , 4. Write the formulas of the (a) primary, (6) secondary, and (c) tertiary calcium salts of orthophosphoric acid. 5. Write an equation for the reaction by which magnesium-ammonium- phosphate is formed when magnesium chloride, ammonia, and ammonium chloride are added to a solution of disodium phosphate. 6. Write an equation for the reaction by which PH 3 and H 3 PO 4 are formed as the result of heating H 3 PO 3 . 7. What principle is a guide in coming to a conclusion as to whether or not the compound having the formula P 4 S 3 will burn? Write an equation for the reaction which takes place when it burns. 420 INORGANIC CHEMISTRY FOR COLLEGES 8. Describe a simple experiment which would show that the reaction between antimony chloride and water is a reversible one. 9. An aqueous solution of I 2 in the presence of NaHCO 3 will oxidize As 2 O 3 to As 2 O 5 . (a) Write an equation for the reaction. (6) How could the reaction be used to determine 12 quantitatively. 10. Name an antidote for each of the following: (a) I 2 , (6) BaCl 2 , (c) HgCl 2 , (d) AgN0 3 . 11. Write equations for all the reactions involved in the Marsh test for arsenic. 12. Write equations for the reactions which take place when (a) Sb 2 O 3 is dissolved in HC1 and (6) the solution is treated with H 2 S; when (c) the sulphide formed is dissolved in (NH 4 ) 2 S 3 and (d) the resulting solution treated with HC1. 13. Starting with Ca 3 (PO 4 ) 2 , write equations for reactions by which the following can be prepared: (a) H 3 PO 4 , (6) H 3 PO 2 , (c) POC1 3 , (d) PH 4 I, (e) Na 4 P 2 O 7 . 14. How could you distinguish from each other the following: (a) Na 2 HPO 4 and Na 2 HAsO 4 , (6) H 3 PO 4 and H 3 PO 3 , (c) AsH 3 and SbH 3 , (d) BiOCl and SbOCl, (e) Bi(OH) 3 and Sb(OH) 3 ? 15. One type of fiiction matches is made by first dipping small wooden sticks in paraffin and then coating the ends with a mixture of potassium chlorate, glue, powdered flint, and clay. On the end of the tip is placed a mixture of red phosphorus (or P 4 S 3 ), potassium chlorate, glue, and clay. The heads of safety matches may contain antimony trisulphide, in addition to the above; the red phosphorus and powdered flint are placed on the box. Explain why each material is used. CHAPTER XXIX SOME IMPORTANT ORGANIC COMPOUNDS 496. The compounds of carbon which were described in Chap- ter XV are all simple in composition and can readily be obtained from mineral sources. They exhibit properties, in the main, like those of the analogous substances derived from the other elements. A large number of compounds of carbon are pro- duced as the result of the life processes in plants and animals and from these many others have been made. Carbon compounds of this kind are considered in detail in a separate branch of chemistry, called organic chemistry, because the principles underlying their chemical behavior are quite different from those arrived at from the study of inorganic substances. The complexity in composition of the compounds of carbon results from the fact that the atoms of the element can unite with each other and thus form molecules, which may contain as many as 60 atoms of carbon. The formulas of some of the simpler hydro- carbons, which are compounds of hydrogen and carbon, will illus- trate this fact. The graphic formulas of methane, CH4, ethane, C2He, and propane, CsHg, are, respectively, as follows: H H H H H H H C H, H C C H, and H C C C H i U ii A The study of organic compounds has centered largely around the determination of the way in which the atoms are joined together in the molecules of these substances. When a number of atoms are present it is possible for two substances to have the same compo- sition but entirely different physical and chemical properties. For example, there are two compounds of the formula CoHeO; it has been shown that in one the atoms are arranged as represented by 421 422 INORGANIC CHEMISTRY FOR COLLEGES formula 1 below, and in the other the arrangement is that indicated in formula 2: H H H H I I I (1) H C C O H (2) H ( H 1 i i i The first is the graphic formula of grain alcohol, a liquid which boils at 76, and the second that of methyl ether, which is a gas. Such compounds are said to be isomeric, because they contain the same number of the several atoms. Only a few of the more important compounds will be mentioned below. 497. Natural Gas and Petroleum. The gas that exists in the earth and is obtained by boring wells consists chiefly of methane, CH4, with which small amounts of hydrogen and ethane, C2He, are mixed. It is usually found in the localities which furnish petroleum. Petroleum occurs in large quantities in the United States in Pennsylvania, Ohio, California, and Texas. It is obtained also in Mexico, certain parts of Central Europe, Japan, and India. Petroleum is a thick, greenish oil which consists essentially of a mixture of a large number of hydrocarbons. The crude material is distilled and in this way separated into products which are sold under trade names for specific purposes. Gasoline, which is used as a motor fuel, is variable in composition, but consists largely of hydrocarbons which boil between 70 and 120. That obtained from most American petroleums contains the hydrocarbons hexane, CeHi4, heptane, CjHie, and octane, CsHig. Kerosene consists of hydrocarbons which boil at higher temperatures and contain from 10 to 16 carbon atoms. Vaseline contains the hydrocarbons of the formula C22H46 and C23H48. Paraffin is a mixture of hydrocarbons of high molecular weight, which are crystalline at room temperature. The hydrocarbons mentioned above belong to what is called the paraffin series; they are characterized by great inertness to chemical reagents. 498. Unsaturated Hydrocarbons. A large number of hydro- carbons are known which contain a smaller proportion of hydrogen than the corresponding compounds in the paraffin series; thus, the formula of ethane is CoHe, and that of ethylene CoH4. The SOME IMPORTANT ORGANIC COMPOUNDS 423 compounds related to ethylene react readily with chlorine, bro- mine, sulphuric acid, and other reagents. In order to indicate this fact, the graphic formula of ethylene is written as given below in the following equation, which represents the union of the hydrocarbon with bromine : H H H H II II C= C + Br 2 = Br C C Br II II H H H H Ethylene is said to be unsaturated because it can unite with other substances by direct addition. This behavior results from the fact that in the hydrocarbon two carbon atoms are in combina- tion with four univalent atoms, whereas the two carbon atoms can be united and hold in combination six univalent atoms. Unsaturated hydrocarbons are present in coal gas, and as they burn with a luminous flame they are the so-called illuminants. Acetylene, (217), is still more highly unsaturated; it has the formula HC^CH. 499. Carbohydrates. Plants are composed largely of com- pounds of carbon, hydrogen, and oxygen, that are called carbo- hydrates. Cellulose, (CeHioOs)*, is the chief constituent of the woody fiber of trees and other plants. It exists in the pure con- dition in cotton and linen. Paper is made from a pulp prepared by beating cotton or linen rags with water. The mixture is placed on a perforated screen and, after the water has drained off, the thin layer of moist pulp is compressed and dried by being passed through heated rollers. Paper is also made in a similar way from wood, after the latter in the form of small chips has been heated under pressure with a solution of calcium bisulphite to dissolve the gums and substances other than cellulose in the wood. In order to prevent the spreading of ink on paper the surface is sized by covering it with a thin layer of gelatine or rosin. Starch, (CeHioOs)!,, is obtained from corn, wheat, potatoes, and other plants. It consists of minute granules which have a characteristic appearance when examined with a microscope. The granules are covered with a very thin coating of a form of cellulose. When starch is heated with water the coatings burst, and when the mixture is cooled a jelly is obtained. A part of the 424 INORGANIC CHEMISTRY FOR COLLEGES starch passes into colloidal solution (534). When a dilute solu- tion of iodine comes into contact with starch that has been heated with water, a characteristic blue color is formed, which serves as a test for either iodine or starch (412). Wheat flour consists of starch and a small proportion of gluten, which is a mixture of certain nitrogenous compounds that are known as proteins. Gluten is essential in bread-making, because it is sticky and, there- fore, prevents the escape of the gas produced when dough rises. Glucose, CeH^Oe, is present in grapes, and for this reason is sometimes called grape sugar. It is also present in honey, and is formed when starch is heated with water; if a trace of an acid is present the change takes place much more quickly. Glucose is made in this way from starch, and after the neutralization of the acid is sold in the form of a syrup. Sucrose, Ci2H22On, is obtained from the sugar cane or sugar beets; it is commonly called cane sugar. When sugar is heated with water in the presence of a small amount of an acid or any substance which furnishes hydrogen ions, it hydrolyzes and glu- cose and fructose, another sugar present in honey and certain fruits, are formed : C 12 H 22 On + H 2 = C 6 Hi 2 6 + C 6 H 12 6 The two simpler sugars contain the same number of carbon, hydrogen, and oxygen atoms; they differ in the structure of their molecules, that is, in the way in which the atoms are joined to one another. This change of sugar into glucose and fructose is brought about in candy-making by heating sugar with cream of tartar (501), which is an acid salt. The mixture formed does not crys- tallize readily and, therefore, gives the candy a smooth texture. Lactose, Ci 2 H 22 On, is a sugar which is obtained from milk; the latter contains, in addition to lactose, fat and casein, which is a protein. 500. Alcohols. When yeast is put into aqueous solutions of most sugars fermentation takes place, and alcohol and carbon dioxide are formed. The reaction in the case of glucose is as fol- lows: C 6 Hi2O 6 = 2C 2 H 6 O + 2CO 2 Yeast contains substances which change the more complex car- bohydrates into glucose and, as a consequence, fermentation takes SOME IMPORTANT ORGANIC COMPOUNDS 425 place when yeast is mixed with moist flour. It is in this way that the dough for bread-making is prepared. Alcohol is obtained by fermentation of the carbohydrates in grains. On account of the fact that it reacts with strong acids to form compounds as the result of the replacement of a hydrogen and an oxygen atom, its formula is usually written C2HsOH. The radical C2H.5 occurs in a large number of compounds and has been given the name ethyl; for this reason grain alcohol is called ethyl alcohol. Alcohol boils at 76 and burns with a light blue flame. Methyl alcohol, CHaOH, is formed when wood is heated to make charcoal; for this reason it is often called wood alcohol. It is obtained by condensing the vapors given off and distilling the resulting liquid. Methyl alcohol is very poisonous and produces blindness. When it burns, carbon dioxide and water are formed, but if an insufficient amount of air is used partial oxidation takes place and formaldehyde, CEbO, is produced. The latter is a gas which dissolves in water. A 40 per cent solution of formaldehyde in water is sold under the name formaline, and is used as a disin- fectant. 501. Organic Acids. The acids derived from carbon have the O grouping of elements represented by the formula C OH (or COOH), which is called the carboxyl group. The formula of O carbonic acid represented in this way is HO C OH. Acetic acid has the formula CHs-COOH. It is obtained along with methyl alcohol when wood is distilled. The acid is also formed as the result of fermentation when fruit juices are exposed to the air, for the latter under ordinary circumstances contains micro- organisms which convert the sugars present into acetic acid. It is in this way that vinegar is formed from the juice of apples or from wine. The organisms which bring about the change are present in what is called mother of vinegar. Lactic acid, C2H5O-COOH, is formed when the sugar in milk ferments; it is present in sour milk. The salts of a large number of acids are present in fruits. Tartaric acid is a dibasic acid obtained from grapes; it has the 426 INORGANIC CHEMISTRY FOR COLLEGES formula C 2 H40 2 (COOH)2 (or H 2 -CJ^Oe); cream of tartar is the acid potassium salt. Citric acid, CsHsCXCOOH^, is obtained from lemons, benzoic acid, CeHs-COOH, from cranberries, and oxalic acid, (COOH)2 (or H2C2O4), from rhubarb and other plants. Oxalic acid forms a dihydrate, H2C2C>4,2H2O, when crystallized from water. It is manufactured by heating sawdust with sodium hydroxide and a small amount of water. From the sodium salt formed in this way, the free acid is obtained by treatment with sulphuric acid. 502. Esters. The products formed as the result of the inter- action of alcohols and acids are called esters; ethyl acetate, for example, is the ester formed from ethyl alcohol and acetic acid; it has the formula CH3-COOC 2 H 5 . CHs-COOH + C 2 H 5 OH CH 3 -COOC 2 H 5 + H 2 O The reaction indicated is reversible, for when an ester is heated with water it is converted into an acid and an alcohol. The pleasant odor of many plants results from the presence of esters. Fats are mixtures of esters of acids of high molecular weight. The alcohol obtained from them when they are heated with water or an alkali is glycerine, CsH^OH^. This reaction is what occurs in soap making. Among the esters in beef fat are the glycerine esters of palmitic acid, C stearic acid, CiyHss-COOH, and oleic acid, C which is an unsaturated acid. When the ester of stearic acid, which has the formula (CiyH^-COO^CsHs, is heated with a solution of sodium hydroxide, sodium stearate, CiyHss-COONa, and glycerine are formed. The mixture of salts produced in this way from fats is soap. The behavior of soap with water and its use as a cleansing agent are described later (628, 629). Fats from different sources differ from one another in the esters present and their relative amounts. Butter is characterized by containing about 5 per cent of the ester of butyric acid, CsHT-COOH. The liquid fats, such as olive oil, contain a high percentage of esters derived from an unsaturated acid. Stearin, (CiTHss'COO^CsHs, which occurs in beef tallow, is a solid, and olein, (Ci7H33-COO)3C3H 5 , which is present in edible oils, is a liquid at ordinary temperatures. SOME IMPORTANT ORGANIC COMPOUNDS 427 Cellulose, which is an alcohol, forms an ester when treated with nitric acid; cellulose trinitrate is used as an explosive (364). A mixture of camphor and the nitrates of cellulose which contain less nitrogen than the trinitrate is called celluloid. The nitrate made from glycerine, CsH^NOa^, is the important constituent of dynamite (364). CHAPTER XXX SILICON AND BORON. THE ACID-FORMING ELEMENTS AND THE PERIODIC CLASSIFICATION 503. Silicon is the second member of the carbon family, which is in the fourth group in the periodic classification. It resembles carbon closely in many chemical properties, but is more active toward electronegative elements; it unites readily with chlorine, for example, whereas carbon does not react directly with the halogen. Its greater activity is also shown by the fact that when sodium carbonate is heated to a high temperature with silicon, sodium silicate is formed and carbon is set free. Silicon has the valence 4 in its compounds with either positive or negative elements; the only compound in which it does not show this valence is silicon monoxide, SiO; it always acts as an acid-forming element. The consideration of tin and lead, two important elements in the fourth group of the periodic classifica- tion, will be deferred until later on account of the fact that their most characteristic properties are those of metals. 504. Occurrence. Next to oxygen, silicon is the most abundant element in the earth's crust, of which it constitutes about 26 per cent. It is never found free, but in combination with oxygen as silicon dioxide, an impure variety of which is sand, or in the form of silicates, which are salts derived from the acid EbSiOa, and are the chief constituents of rocks and clays. The soil is a mixture of sand, silicates, and the complex organic compounds formed as the result of the decomposition of dead vegetable material. The silicates that occur in nature and those made artificially have extensive uses on account of their physical properties. 505. Preparation. Silicon is obtained by heating a mixture of sand and carbon in an electric furnace (218); carbon monoxide and silicon are the chief products of the reaction, but some silicon monoxide is produced in the parts of the furnace which are at a 428 SILICON AND BORON 429 lower temperature than the central core through which the major part of the current passes. Ferrosilicon, an alloy of iron and sili- con, is prepared on the large scale by a similar process from a mix- ture of oxide of iron, sand, and carbon; it is used in the manu- facture of steel. Silicon is an important constituent of the crude iron obtained from blast-furnaces (pig-iron), and has a marked effect on its properties. In the laboratory, silicon is prepared by igniting a mixture of powdered silicon dioxide and magnesium. It is necessary only to start the reaction, which is a vigorous one: SiO 2 + 2Mg = 2MgO + Si The silicon is obtained by treating the mixture with an acid, which dissolves the magnesium oxide and the small amount of magnesium silicide, Mg 2 Si, present. Silicon can be obtained in a crystalline condition by heating it with molten zinc, and after the mass has solidified, dissolving out the metal with an acid. 506. Properties. Amorphous silicon, which is a brown powder, is more active than the crystalline variety; at ordinary tempera- tures it unites with fluorine and, when heated, with chlorine, bro- mine, oxygen, and sulphur. At high temperatures it combines with nitrogen, and at the temperature of the electric furnace with carbon and boron. It is slowly oxidized by aqua regia to silicic acid, and is dissolved by a mixture of hydrofluoric and nitric acids. Silicon forms silicides with the more active metals. Silicon reacts with solutions of the alkalies: 2NaOH + Si + H 2 O = Na 2 SiO 3 + 2H 2 The reaction is utilized in making hydrogen for filling balloons. Crystalline silicon is obtained in the form of black needles, which are harder than glass and will, therefore, scratch it. In order to bring it into reaction a higher temperature is required than is the case with the amorphous element. 507. Silicon Dioxide. Silicon dioxide, which is also called silica, exists in nature widely distributed in a more or less pure condition as sand, which is ordinarily colored as the result of the presence of compounds containing iron. Sandstone is made up of sand grains cemented together by amorphous silica, clay, and other substances; it is extensively used for building purposes. 430 INORGANIC CHEMISTRY FOR COLLEGES Quartz or rock crystal is the name given to a crystalline variety of silicon dioxide, which occurs as clear, colorless, six-sided prisms terminated by pyramids. Many varieties of quartz are found that are colored as the result of the presence of small amounts of impurities; among these are amethyst, which contains traces of manganese and iron, and yellow, rose, and smoky quartz, and cat's eye. If the silica is not crystalline and is colored brown or red by iron oxide, it is known as agate, onyx, or jasper. Opal and flint are varieties of hydrated silica. Infusorial or diatomaceous earth, which is also called Tripoli powder, is a form of silica made up of the skeletons of minute organisms. On account of its physical condition it is used as a polishing powder, and on account of its absorbent qualities for removing the coloring matter from oils. Rock crystal is used in making special kinds of chemical appa- ratus required when reactions are to be carried out at very high temperatures or when a vessel must be used that is not attacked by acids. Since quartz melts at 1600 the apparatus must be made with the aid of the oxy-hydrogen flame or in an electric fur- nace. Apparatus made from rock crystal is transparent and resembles glass in appearance; it possesses one property that is very valuable it expands to such a slight degree when heated that it can undergo great changes in temperature without crack- ing; a quartz vessel can be heated red-hot and plunged into cold water without breaking. Apparatus made from other varieties of silicon dioxide are milky in appearance, but on account of their lower cost are used in place of that prepared from rock crystal when transparency is not required. It has recently been found possible to construct silica apparatus of such a size that it can be used in chemical operations carried out on the industrial scale. Quartz is much more transparent to ultra-violet light than glass and is used in place of the latter in experimental work with this form of energy. 508. Silicic Acid and Silicates. Although silicon dioxide does not react with water to form an acid, it possesses the most char- acteristic property of an acid anhydride it dissolves in solutions of the alkalies to form salts, which are known as silicates. As the reaction takes place very slowly, sodium silicate, which has extensive uses, is made by fusing sand with sodium hydroxide or sodium carbonate. When an acid is added to a solution of sodium SILICON AND BORON 431 silicate, silicic acid is formed and appears as a gelatinous precipitate. The composition of the precipitate is approximately that repre- sented by the formula of orthosilicic acid, H4SiO4. When it is dried it slowly loses water until silicon dioxide is formed. A large number of salts are known derived from metasilicic acid, H 2 SiO 3 , which resembles carbonic acid in composition. 509. Many minerals are derived from what are called poly- silicic acids, which are formed as the result of the partial dehy- dration of two or more molecules of orthosilicic acid; they may be considered as made up of silicon dioxide in combination with water in different proportions, the relation between the acids being similar to that already seen in the case of the phosphoric acids. If 1 molecule of orthosilicic acid loses 1 molecule of water, it is transformed into metasilicic acid, H 4 SiO 4 - H 2 O = H 2 SiO 3 (H 2 O,SiO 2 ) Polysilicic acids are derived from two or more molecules of the ortho acid; the more important of these contain 2 and 3 atoms of silicon. The salts of two disilicic acids are known ; one acid is formed from 2 molecules of the ortho acid by the loss of 1 molecule of water, 2H 4 SiO 4 - H 2 O = H 6 Si 2 O 7 (3H 2 O,2SiO 2 ) and the other by the loss of 3 molecules, 2H 4 SiO 4 - 3H 2 O = H 2 Si 2 O 5 (H 2 O,2SiO 2 ) Trisilicates are known which are derived from an acid formed as the result of the loss of 4 molecules of water from 3 molecules of orthosilicic acid, 3H 4 SiO 4 - 4H 2 O = H 4 Si 3 O 8 (2H 2 O,3SiO 2 ) 510. A large number of minerals are salts of these four silicic acids, but only a few will be mentioned. Garnet, which is some- times used as a jewel, is an orthosilicate of magnesium and alumin- ium in which more or less of the magnesium is replaced by calcium. To express this latter fact the symbols (Mg,Ca) are written in the formula, which is (Mg,Ca) 3 Al 2 (SiO 4 ) 3 . Mica is a familiar mineral that is used for a number of purposes on account of its peculiar physical structure which allows it to be separated into thin sheets; it is also an orthosilicate and has the formula 432 INORGANIC CHEMISTRY FOR COLLEGES (K,Na)H2Al3 (8104)3: it is extensively used in the construction of electrical apparatus because it does not conduct electricity and resists fairly high temperatures. Kaolin, clay, is a hydrated aluminium orthosilicate, H 2 Al 2 (SiO 4 ) 2 ,H 2 O. Serpentine is an example of a disilicate and has the formula Mg 3 Si 2 O7,2H 2 O. Feldspar, or orthoclase, which plays an important part in soil- formation, and is used in making porcelain and other table- ware, is a trisilicate and has the formula KAlSisOg. A large number of metasilicates are important minerals. Among these are talc (soapstone), H 2 Mg 3 (810)3)4, and asbestos, MgsCa (8163)4. The minerals that occur as silicates are the chief constituents of many igneous rocks, which are usually mechanical mixtures of two or more minerals. Granite, for example, which is a very widely distributed rock, contains feldspar, mica, and quartz. 511. Sodium silicate is manufactured on a large scale, as it has important industrial applications; it is known as water-glass on account of the fact that it is soluble in water, and is obtained as a transparent amorphous mass resembling glass when its solu- tion is evaporated to dry ness. It is made by fusing together for eight to ten hours at a high temperature a mixture of quartz or infusorial earth and caustic soda, sodium carbonate, or sodium sul- phate. The product is a glass which is powdered and heated with water under pressure until the soluble material is dissolved. The solution is then evaporated until it has the specific gravity 1.7. The composition of the silicate formed is determined by the proportions of the ingredients used; it approximates that of the formula Na 2 Si4Og (Na 2 Si03,3SiO 2 ). It is used as a fixative for pigments in calico printing, for rendering cloth and paper non- inflammable, for preserving eggs, in cement mixtures for glass, wood, and leather, as a preservative for timber and porous stone, in the manufacture of artificial stone, and as a size for paper and fabrics. Sodium silicate is mixed with soap in the preparation of certain varieties of the latter used for laundry work, because it serves the valuable purpose of softening hard-water. All the silicates are insoluble except those of the alkali metals. Since silicic acid is a very weak acid, its soluble salts are hydro- lyzed by water and show an alkaline reaction. The use of sodium silicate as a paint-remover is based on this fact; the alkali formed SILICON AND BORON 433 as the result of the hydrolysis of the salt attacks the solidified oil in the paint and converts it into substances soluble in water. 512. Test for Silicates. When an acid is added to a solution of a silicate, silicic acid is precipitated in a gelatinous condition. To complete the identification of the latter it is filtered off, treated with hydrofluoric acid, and evaporated to dryness; silicic acid leaves no residue, as it is converted by hydrofluoric acid into silicon fluoride, which is volatile. If the substance thought to be a sili- cate is insoluble, it is fused at red-heat with sodium carbonate, which converts insoluble silicates into soluble sodium silicate, and this is tested as just described. 513. Hydrides of Silicon. Silicon forms hydrides of the com- position SiH 4 , Si 2 He, and Si2H2. They are all gases which burn in air. The tetrahydride can be prepared by treating magnesium silicide with hydrochloric acid : Mg 2 Si + 4HC1 = 2MgCl 2 + SiH 4 When prepared in this way it contains an impurity which renders it spontaneously inflammable; it is readily decomposed by heat into its constituents. 514. Halides of Silicon. Silicon tetrachloride can be pre- pared by the direct union of the elements, or if silicon is not avail- able, by the action of chlorine and carbon at a high temperature on silicon dioxide: SiO 2 + 2C + 2C1 2 = SiCl 4 + 2CO Neither carbon nor chlorine alone will react with silica at the temperature which can be used to prepare the chloride in this way. In the simultaneous action of carbon and chlorine the affinity of the former for oxygen and the latter for silicon come into play at the same time and the change takes place. Silicon tetrachloride is a colorless liquid, which boils at 59 and fumes in the air as the result of the hydrolysis brought about by the moisture present: SiCl 4 + 4H 2 O = Si(OH) 4 + 4HC1 When the vapor of the chloride comes in contact with air contain- ing ammonia a very dense cloud is formed. The reaction was 434 INORGANIC CHEMISTRY FOR COLLEGES utilized in the recent war in producing smoke clouds to hide vessels at sea. Silicon tetrafluoride (426) can be prepared by heating together a mixture of calcium fluoride, sand, and concentrated sulphuric acid; hydrofluoric acid is first formed, and then reacts with the silicon dioxide to produce silicon fluoride and water. Silicon tetrafluoride is a gas, which reacts vigorously with water: SiF 4 + 2H 2 O ^ Si(OH) 4 + 2H 2 F 2 The hydrofluoric acid produced unites with some of the silicon fluoride present and forms hydrofluosilicic acid, SiF 4 + H 2 F 2 = H 2 SiF 6 which can be considered as derived from silicic acid, H 2 Si03, as the result of the replacement of 3 oxygen atoms by 6 fluorine atoms. Hydrofluosilicic acid is stable in solution only. When a solu- tion is evaporated, silicon fluoride is given off. The acid can be used for testing for potassium, as its salt containing this metal is difficultly soluble in water. BORON 515. Boron is the first member in the third group in the periodic classification of the elements. It exhibits in most of its compounds the properties of an acid-forming element, but it forms an acid sul- phate and a phosphate, which are readily hydrolyzed by water. The second member of the group, aluminium, is a well-charac- terized metal. The compounds of boron which occur in nature and those which have important uses are derivatives of boric acid. The element has a strong affinity for oxygen and in the course of the formation of the earth united with this element; since the oxide is an acid anhydride, salts were produced as the result of combina- tion taking place between it and metallic oxides. 516. Occurrence. Boric acid, HaBOs, is a very weak acid, and like other weak oxygen acids it forms salts derived from acids the molecules of which contain a number of the atoms of the acid- forming element. The polyboric acids bear a relation to normal boric acid, HsBOs, similar to that which exists between silicic acid, H 4 Si0 4 , and the polysilicic acids. SILICON AND BORON 435 Three important minerals containing boron are borax, (Na 2 O,2B 2 O 3 ), which is a salt of tetraboric acid; cole- manite, Ca 2 B 6 Oii,5H 2 O (2CaO,3B 2 O 3 ,5H 2 O) ; and boracite Mg 3 B 8 Oi5,MgCl 2 (3MgO,4B 2 O 3 ,MgCl 2 ). Since the salts of boric acid are hydrolyzed by water, the free acid is formed in nature in certain hot springs; those of Tuscany are used as a commercial source of the acid. 517. Preparation. The method used to isolate boron is analo- gous to that employed to obtain free silicon; its oxide is heated with magnesium, and the resulting mass treated with an acid to dissolve the magnesium oxide formed. The product is amor- phous boron, which can be obtained as crystals by dissolving it in molten aluminium and removing the latter, after cooling, by means of hydrochloric acid. Aluminium can also be used to reduce boron oxide, and if an excess of the metal is used the product obtained directly is crystalline boron. 518. Properties. Amorphous boron is a black powder, which can be fused at the temperature obtained in an electric furnace. It is brittle, very hard, and a poor conductor of heat and elec- tricity at ordinary temperatures. With rise in temperature its physical properties change rapidly; at 400 its electrical con- ductivity is over two million times that shown at room-tempera- ture, and it approaches a metal in electrical properties. At high temperatures boron is a very active element, which resembles silicon and carbon in this respect, but it is more active than these elements, for it will decompose both silicon dioxide and carbon monoxide. When boron is heated with nitrogen, it forms a nitride, BN, which is converted by steam into boric acid, H 3 BO 3 , and ammonia. At the temperature of the electric furnace boron and carbon form a carbide, BeC, which is characterized by its great hardness, as it stands next to the diamond in this respect. Free boron is added in small quantities to copper when the latter is to be used for making castings because it gives a product which is free from blow-holes. Boron dissolves in molten potassium hydroxide: 2B + 6KOH = 2K 3 BO 3 + 3H 2 519. Boric Acid. The acid is found free in nature as has been indicated. In Tuscany, ponds are formed around the fumeroles 436 INORGANIC CHEMISTRY FOR COLLEGES from which steam issues carrying boric acid (suffioni) ; the steam is condensed in the water, and a solution obtained which contains about 2 per cent of the acid. This is evaporated by the steam from the fumeroles and crystalline boric acid is obtained. Boric acid, H 3 BO 3 , is made in California and in Chile, by heating calcium borate (colemanite) suspended in water, through which sulphur dioxide is passed; the sulphurous acid formed con- verts the calcium borate into boric acid and calcium bisulphite. Boric acid crystallizes from water in pearly white, thin plates. It is soluble in 25 parts of water at 19 and in 3 parts at 100. When heated, it loses water; at 140 tetraboric acid, H^B^y, is formed and at red heat it is converted into its anhydride, B9Os, which is non- volatile at high temperatures. Both tetraboric acid and boron trioxide react slowly with water to form boric acid. Boric acid is a very weak acid; its aqueous solution barely affects litmus paper; it is slowly volatile with steam. Boric acid is used in the preparation of borax, and in making enamels and glazes for pottery and in certain kinds of glass. Its solution has antiseptic properties and is used as an eye-wash and for other purposes in medicine under the name boracic acid. It is also used for preserving meat, fish, oysters, and milk, although its use for this purpose is prohibited by law in certain States. 520. Berates. The only salt derived from boric acid which is used extensively is borax, sodium tetraborate, ^26467 . It is obtainable commercially as the pentahydrate, ^26467, 5H2O, or as the decahydrate, Na2B4O7,10H2O. Borax is found native in Thibet, Ceylon, and California. It was formerly prepared in large quantities from the water of Borax Lake, California, by evaporating it to crystallization. It is now obtained from dry lake beds in the Death Valley region, where the surface of the earth is covered with a crust composed of borax and the sulphate, chloride, and carbonate of sodium. Most of the borax pro- duced in California is made from colemanite or from ulexite, NaCaBsOojSH^O. The mineral in either case is roasted, and as the result of the loss of water falls to a powder; this is then boiled with a solution of sodium carbonate and sodium bicarbonate; the calcium is precipitated as carbonate, and borax is obtained from the solution by evaporation. The decahydrate of borax is obtained when the salt crystallizes SILICON AND BORON 437 from solution at 27 or below; it forms large efflorescent prisms, which melt in the water of hydra tion when heated, and swell as the latter is given off, forming a spongy mass that at higher tempera- tures melts to a clear glass. If a solution of borax is concentrated so that crystals are formed at 56 or higher, the pentahydrate is obtained in the form of octahedral crystals, which are permanent in dry air; it fuses without swelling and is, therefore, preferred to the pentahydrate when used as a flux. Borax is hydrolyzed to the extent of about one-half per cent in tenth-normal solution, and shows, accordingly, an alkaline reac- tion. Some of the uses to which borax is put depend upon this fact. It is employed, for example, in ungumming raw silk; the small amount of alkali formed in the hydrolysis serves to dissolve the gum from the fiber, which would be seriously affected if a solution of sodium hydroxide were used. The use of borax in soap is based in part upon the fact that the alkali formed from it markedly assists in converting fat and oily substances into such a finely divided condition (an emulsion) that they can be readily removed by water. If soap is to be used with hard water the presence of borax is advantageous since the soluble calcium salts present in the water are precipitated as calcium borate and do not, therefore, interfere with the action of the soap. Borax is used for preserving meat, as a mordant in dyeing, in medicine and pharmacy, as a flux, and in the manufacture of enamels and glazes for metal ware and pottery. 521. A flux is a substance which converts compounds infusible at a certain temperature into others which melt at this tem- perature. Fluxes are generally used when reactions are carried out with refractory substances at high temperatures, on account of the fact that the chemical reaction takes place only when the substances involved are in intimate contact a condition difficult to bring about in the case of solids. The use of borax as a flux is based upon the fact that it contains an excess of boric anhydride over that required to form sodium metaborate, NaB(>2. The relation in com- position between the two salts is clearly seen by comparing their two formulas written in such a way as to indicate the amounts of boron oxide present; sodium metaborate is Na20,B203, and borax is Na2O,2B2C>3. When the oxide of a metal is heated with fused borax, it dissolves as the result of its union with the excess of boric 438 INORGANIC CHEMISTRY FOR COLLEGES anhydride to form a metaborate. For example, copper oxide and borax under these conditions form a mixture of sodium metaborate and copper metaborate, which is fusible: CuO + Na 2 O,2B 2 O 3 = CuO,B 2 O 3 + Na 2 O,B 2 O 3 A similar reaction is made use of in welding together two pieces of iron; the parts to be united are covered with borax, heated to redness, and hammered; the oxide on the surface of the metal does not melt at the temperature used, and is converted into iron metaborate, which mixes with the molten borax and is thus expelled from the point of contact of the two pieces of metal; they can now be welded together in the soft condition of the iron at the tem- perature used. Borax is used for the same reason in soldering. The property shown by borax of dissolving oxides of metals is utilized in qualitative chemical analysis in the so-called borax- bead test. A bit of borax is placed in a small loop in the end of a platinum wire and melted down to a clear glass in the Bunsen flame. While hot it is brought into contact with the substance to be tested and returned to the flame. At the high tempera- ture, salts are decomposed and the resulting oxides dissolve in the molten borax as metaborates. Many of these salts possess char- acteristic colors, and the metal present can, therefore, be iden- tified. Further, in some cases when a metal forms salts in which it can exhibit different valencies, the color of the salts is different when it is in one condition from that when in the other. If the bead is held in the oxidizing part of the Bunsen flame near the tip, the metaborate formed will be derived from the metal in its higher state of oxidation; if it is held in the reducing flame, near the top of the blue inner cone, the salt will be derived from the metal in its lower state of oxidation. 522. Microcosmic salt, sodium-ammonium phosphate (462), can be used instead of borax in the bead test. This salt is converted into sodium metaphosphate, NaPOs, when fused, and when heated with oxides of metals dissolves them with the formation of orthophosphates; for example, with copper oxide the following reaction takes place: CuO + NaP0 3 = CuNaP0 4 SILICON AND BORON 439 523. Tests for Boric Acid and Borates. The most important test for these substances is based on the fact that certain organic salts of boric acid impart a green color to the flame produced when they burn. The test is carried out by treating the substance to be examined with alcohol and concentrated sulphuric acid; when a flame is applied to the mixture, the alcohol and the ethyl borate, (C2H5)3BO3, formed burn, and a green flame is produced. A second test which is convenient is carried out by first moist- ening a piece of turmeric paper with a solution of the borate con- taining a little hydrochloric acid, and then drying the paper by placing it on a flask containing boiling water. The paper, which was originally yellow, turns to a rose pink on drying if boric acid is present, and if put into a dilute solution of ammonia the pink color changes to bluish black. 524. Other Compounds of Boron. Boron forms two hydrides, B4Hio, and B2He, a trichloride, and a trifluoride. Boron trifluoride can be prepared by a reaction analogous to that used to make silicon fluoride, and, like the latter compound, reacts with water to form a complex acid; a number of salts of hydrofluoboric acid, HBF4, are known. THE ACID-FORMING ELEMENTS AND THE PERIODIC LAW 525. The elements described up to this point are the more important ones possessing acid-forming properties. From time to time the relation between their physical and chemical properties and their position in the periodic classification of the elements has been emphasized, and the effect of the valence of an element on the properties of the compounds containing it has been pointed out. It will be well worth while to bring together here the more important generalizations, as these will help materially to fix in the mind facts already learned, and will be the best preparation for the study of the chemistry of the metals and their compounds. 526. Relationships in a Chemical Family. Physical Proper- ties. With increasing atomic weight the densities of the elements and their compounds increase, and their boiling-points, melting- points, and critical temperatures also increase. The number of atoms in the molecules of the elements in the gaseous condition, either remains the same, as in the case of the halogens, or decreases as seen in the phosphorus family. 440 INORGANIC CHEMISTRY FOR COLLEGES Chemical Properties. The stability toward heat of the molecule of the element in the gaseous condition decreases with increasing atomic weight; this is shown in the cases of chlorine and iodine and phosphorus and arsenic. The activity of the elements as measured by the heat of forma- tion of the hydrides and other compounds containing electro- positive elements decreases with increasing atomic weight, a fact which is associated with the decreasing stability of the hydrides with increasing atomic weights. The heats of formation of the hydrides from the elements in the gaseous condition are positive in most cases. As the affinity of the elements for hydrogen decreases their electro-positive properties increase. Since in the salts of the acids containing oxygen the element is functioning in its positive character, the element with the higher atomic weight will replace, in general, the one with the lower atomic weight. On the other hand, in the case of the hydrides, in which the elements are electro-negative, the order of replacement is the reverse; chlorine, for example, replaces bromine. The extent of the hydrolysis of the chlorides of the elements decreases with increasing atomic weight, which is another indica- tion of increase in electro-positive nature in this order. The effect of change in valence of an element is great. Most of the oxides of the acid-forming elements unite with water to form acids, the strength of which in the case of a single element is determined by its valence the higher the valence, that is, the larger the proportion of oxygen, the stronger the acid. We have seen, for example, that nitric, sulphuric, and chloric acids are stronger, respectively, than nitrous, sulphurous, and hypochlorous acids. In the case of amphoteric elements like antimony, the hydroxide containing the element in the lower valence is a weaker acid and a stronger base than the hydroxide containing the ele- ment in the higher valence. We shall see later that this important generalization can be extended to include the bases formed from the metallic elements; the lower the valence of the element the stronger the base it forms; sodium hydroxide, NaOH, is a very active base, aluminium hydroxide, A1(OH)3, is a very weak base, which shows some of the properties of acids; iron forms two basic hydroxides, Fe(OH)2 and Fe(OH) 3 , the former being the stronger base. SILICON AND BORON 441 527. Relationships between the Chemical Families. When the elements were taken up systematically according to the periodic classification, the halogen family was first discussed because the electro-negative character of all its members is pronounced, and fluorine is the most electro-negative of all the elements. As we consider the families one after another passing from the right of the table to the left, we discover that the compounds of the ele- ments with oxygen and hydrogen become weaker and weaker acids, and the electro-negative character of the elements decreases. This is, no doubt, associated with their decreasing valence toward oxygen. Since in any family decrease in electro-negative nature follows increase in atomic weight, we have, thus, two factors which tend simultaneously to decrease the electro-negative character of the elements. While in the seventh group all the elements are strongly electro-negative, in the sixth tellurium shows some metallic properties, in the fifth antimony and bismuth, in the fourth all the members below silicon, and in the third even boron, the first member of the group, shows these properties, although its oxide is more acidic than basic. If we draw a line across the periodic table from the square in which boron is placed to that under iodine, we divide the elements roughly into those which are electro-positive and those which are electro-negative; the ele- ments near the diagonal line show amphoteric properties; the most electro-positive and electro-negative elements are furthest from the line. The classification in this way does not include the elements in group which show no chemical properties, and those in group 8, which are distinctly metallic in properties. EXERCISES 1. Write balanced equations for the reactions which take place between the following substances: (a) silicon and aqua regia, (6) water-glass and hydrochloric acid, (c) Colemanite and sodium carbonate when fused together, (d) silicon tetrachloride, water, and ammonia, (e) silicon dioxide, con- centrated sulphuric acid, and calcium fluoride, (/) silicon dioxide and calcium sulphate (fused), (g) lime, sodium sulphate, and silicon dioxide (fused), (h) lead oxide and water-glass (fused), (i) Colemanite, water, and sulphur dioxide, (j) boron nitride and steam, (fc) boracite and hydrochloric acid, (Z) borax and hydrochloric acid, (ra) boric acid and sodium hydroxide, (n) boric acid, alcohol (C 2 H 5 OH), and concentrated sulphuric acid. 2. Give as many examples as you can of the facts summarized in section 526. Tabulate the facts according to the properties listed. CHAPTER XXXI PHYSICAL PROPERTIES OF METALS. ALLOYS The uses to which metals are put are based upon their physical or chemical properties, but as the great majority of these uses depend on physical properties the latter will be discussed briefly, although their consideration rightly falls in physics and engi- neering. 528. Density. The metals vary greatly in density (175); the lightest is lithium, which has the density 0.534 and is, there- fore, about one-half as heavy as water, and the heaviest is osmium (d. 22.5) which is closely related to platinum (d. 21.37) in physical and chemical properties. The so-called light metals, of which sodium, potassium, magnesium, and aluminium are examples, have a density less than 4; iron, lead, tin, silver, etc., are known as heavy metals. In the construction of automobiles and aero- planes the density of the metal used is an important factor, for the power consumed in moving them increases rapidly as the weight increases. For this reason aluminium (d. 2.7) is used for such parts as do not bear a great strain; the metal cannot replace the much heavier steel for many purposes on account of the fact that its tensile strength (see below) is comparatively low. An alloy which contains aluminium, copper, and magnesium is used in making parts for aeroplanes. Magnesium (d. 1.74) is not a strong metal, but forms alloys with other metals which have a low density and a relatively high tensile strength. The table on page 443 gives the densities and some other physical properties of the metals ; it should be consulted in connection with the discussion of these properties which follows. 529. Hardness. The metals vary in hardness, from potas- sium, which can be molded like wax, to chromium, which will cut glass. In the arbitrary scale of hardness which is used to express numerically this property, a number of substances are selected as 442 PHYSICAL PROPERTIES OF METALS. ALLOYS 443 standards (see footnote 4, below) and arranged in such an order that each member of the series can be scratched by the succeeding member. When the hardness of lead is said to be 1.5 the state- ment means that lead will scratch talc, 1, but not rock-salt, 2. PHYSICAL PROPERTIES OF METALS Metal. Density. Melting- point. Boiling- point. Electrical l Eesistance. Heat 2 Conduc- tivity. Tensile 3 Strength. Hard- ness. 4 Aluminium . Antimony. . Barium 2.70(20) 6.62(20) 380 658 630 850 1800 1440 950 2.6-3.0 35.4-45.8 0.48 0.0444 12,590 1,000 2. 3.3 Bismuth Cadmium 9.78(20) 8 65(20) 269 321 1420 778 108 6.2-7.0 0.017 0.22 3,000 2.5 Caesium 1 87(26) 264 670 193 Calcium 1.54(29) 805 10.5 Copper Gold Iron Lead Lithium 8.94 19.32(17) 7.90 11.34 0534 1083 1063 1520 327 186 2310 2200 2450 1525 1400 1.54 2.09 9.7-12.0 18.4-19.6 85 1.00 0.75 0.20 0.084 24,000 20,000 48,000 2,050 2.5-3 2.5-3 4-5 1.5 ]VIagnesium 1 74(5) 651 1120 4 1-50 037 Mercury . . . Nickel 13.595(4) 8.8 -38.85 1452 357.25 2450 94 10.7-12.4 0.015 0.142 54,000 Platinum . . . Potassium 21.37 87(20) 1755 625 2650 712 9.0-15.5 0.166 45,000 4.3 Silver Sodium 10.50 097 961 976 1955 750 1.5-1.7 4.4 1.09 41,000 2.5-3 Strontium . 2.54 900 24.8 Tin 7.3(15) 231.9 2270 9.53-11.4 0.153 4,600 1.5 Tungsten 19 1 3200 70 Zinc . . 7.19 419 940 5.56-6.04 0.265 5,000 2.5 1 The figures given are expressed in millionths of an ohm and are the resistance offered by a cube of the metal 1 centimeter on each edge. 2 The figures given express in calories the quantity of heat transmitted per second through a plate 1 centimeter thick across an area of 1 square centimeter when the difference in temperature between the two sides of the plate is one degree. 3 The figures give approximate values of the number of pounds per square inch required to effect a permanent elongation of the cast metal. 4 The arbitrary scale of hardness used is as follows : 1 talc, 2 rocksalt, 3 calcite, 4 fluorite, 5 apatite, 6 feldspar, 7 quartz, 8 topaz, 9 corundum, 10 diamond, 444 INORGANIC CHEMISTRY FOR COLLEGES The hardness of metals is markedly affected by the presence of other substances in them. While the hardness of pure iron is 4 to 5, steel, which contains carbon, silicon, and other substances, varies in hardness according to its composition from 5 to 8.5; the use of steel in making files is based on its hardness; sandpaper, which is made by covering paper with quartz, hardness 7, is used for polishing wood and other soft objects, but emery paper, made from crushed corundum, hardness 9, is used when iron and other metals are to be abraded. The position of a few common sub- stances in the scale of hardness is of interest; amber is 2.5, asphalt, 1-2, brass 3-4, garnet 7, glass 4.5-6.5, marble 3-4, meerschaum 2-3. 530. Tenacity. The metals and other substances differ in the extent to which they can resist a strain brought to bear on them that tends to bring about a permanent change in their form. In the seventh column of the table on page 443 are given the values of the tensile strength of the more common metals. This property is greatly affected by the purity of the metal and the physical treatment to which it has been subjected, such as casting, rolling, drawing, etc. Pure iron wire, hard drawn, has a tensile strength of from 80,000 to 120,000; after annealing the value is 50,000 to 60,000. The tensile strength of cast iron is from 13,000 to 33,000 and of steel from 80,000 to 460,000. Cast zinc has a tensile strength of from 7,000 to 13,000, whereas the value for the drawn metal is 22,000 to 30,000. The effect of heat treatment on the properties of steel will be considered later in more detail on account of its great importance. 531. Electrical Conductivity. All substances offer more or less resistance to the flow of an electric current through them. With any given substance the resistance is determined by its dimen- sions and the temperature. When these are fixed the resistance is constant and a characteristic property of the substance. Differ- ent substances differ widely in the resistance they offer to the flow of an electric current; in the case of metals it is very small, whereas with the non-metals and most other substances in the solid condi- tion it is very great. The standard of resistance is that offered at by a column of mercury 1 mm. in cross-section and 1.063 meters in length; it is called an ohm. Electrical conductivity is the reverse of electrical resistance and is inversely proportional to it. For example, the resistance PHYSICAL PROPERTIES OF METALS. ALLOYS 445 of silver is about one-sixtieth that of mercury, and, consequently, its conductivity is sixty times that of the latter. The electrical conductivity is denned as the reciprocal of the resistance; it equals numerically 1 divided by the resistance, and is expressed, there- fore, in reciprocal ohms. In the fifth column in the table on page 443 the specific resist- ance of the metals is given, that is, the resistance of a cube of the metal 1 cm. on each edge. Since the numbers are very small they are expressed in millionths of an ohm to avoid large fractions. It is seen that there is a rather wide variation in the values for the several metals; the resistance of silver, copper, gold, and aluminium is comparatively small; magnesium and zinc come next, and platinum, iron, and nickel fall into a group. The resistance of the elements in which metallic properties are not highly developed, like antimony and bismuth, is high. The conductivity of a metal is affected by its physical condi- tion brought about as the result of heat treatment, drawing, casting, etc., and by the presence of even very small amounts of impurities. It is for this reason that copper is refined with the greatest care when it is to be employed in the form of wire for elec- trical construction on account of its low resistance. The figures given in the table for iron, 9.7-12.0, indicate the resistance of the usual grades of commercial iron wire; using the same unit the re- sistance of hard cast iron is 97.8, and that of a soft steel 15.9 and a hard steel 45.7. The high resistance of certain alloys is utilized in the construction of resistance coils for electrical purposes and when it is desired to convert electrical energy into heat to be used in connection with electric stoves, ovens, and furnaces. The heat generated by an electric current is proportional to the resistance. The alloy known as constantan, which contains 60 per cent of copper and 40 per cent of nickel, is used for this purpose; its specific resistance in the unit used in the table is about 49. Platinum can also be used and is of particular value because it does not oxidize at high temperatures, but its Cost prohibits its use except in small apparatus. Nichrome (or chromel) is an alloy of nickel and chromium which is much used in the construction of electrical heating devices. Its specific resist- ance (about 100) is great and it resists the action of air at reason- ably high temperatures. 446 INORGANIC CHEMISTRY FOR COLLEGES It is of interest to compare the resistance of metals with that of the so-called " non-conductors " or dielectrics. The resistance of a centimeter cube of copper is 1.5 millionths of an ohm, while that of glass of the same dimensions is 9 X 10 13 ohms. The values for a few other substances are as follows: quartz crystal 1 X 10 14 , rock salt 9 X 10 16 , rhombic sulphur 70 to 390 X 10 13 , dry wood 5 to 10 X 10 8 , petroleum 2 X 10 16 , and distilled water 0.5 X 10 6 . Solutions of acids, bases, and salts conduct electricity. The following figures express in ohms the resistance of 10 per cent solutions: hydrochloric acid 1.59, sodium hydroxide 3.20, sodium chloride 8.33, zinc chloride 13.75, and copper sulphate 31.2. 532. Heat Conductivity. Metals are the best conductors of heat. In the sixth column of the table of the physical properties of metals will be found the quantity of heat, measured in calories, transmitted per second through a centimeter cube of the metal when the difference in temperature of the two sides of the cube is 1 degree. A comparison of the figures with those for the electrical resistance of the metals brings out clearly the fact that high heat conductivity is associated with low electrical resistance; the good conductors for electricity conduct heat well; high heat con- ductivity is, thus, a characteristic property of metals. The corresponding figures for some non-conductors of heat are as fol- lows: asbestos paper, which is used in making fire-proof paper, theater curtains, etc., 0.00043 calorie; red brick 0.00150; blotting paper 0.00015; cork 0.00072; cotton wool, 0.000043; eiderdown, 0.000011; felt, 0.000087; glass, 0.0012; ice, 0.00396; magnesia, 0.00016-0.00045; quartz, 0.00036; sawdust, 0.00012; silk, 0.00009; dry soil, 0.00033; air, 0.000057; water, 0.0012. 533. The Physical State of Metals. All the metals with the exception of mercury are solids at ordinary temperatures. They can be obtained in the form of crystals by fusing them and pouring off the molten metal after a part has crystallized; or they can be crystallized from a solvent, using for this purpose some other metal with which they do not interact. Most of the metals crystallize in the cubic system. When molten metals solidify they change to a mass of closely interlocked crystals; if such a mass is hammered, rolled, or drawn through a die, the crystals are more or less broken, and due to the pressure exerted on them their faces are forced into closer contact and, perhaps, converted PHYSICAL PROPERTIES OF METALS. ALLOYS 447 into an amorphous condition when they coalesce. If a piece of metal, such as a wire, is heated until it softens, the crystals have an opportunity to grow and rearrange themselves; if now it is cooled rapidly by plunging it into water, it is " frozen " in the condition in which it existed at the higher temperature, and its physical properties are different from what they were before the heat treatment. If a platinum wire which has become brittle is heated to redness and plunged into water it becomes soft and pliable. Most metals are hardened by this treatment. It is evident that pressure and heat treatment have an . effect on the physical prop- erties of metals; this is more notably the case with alloys in which there are several kinds of crystals present in the solid material. The color of metals is determined by their state of division. The color usually associated with a metal is that of the light reflected from its surfaces when it is in the massive condition; silver is white, gold yellow, copper red, etc. In very thin layers metals transmit light, which in certain cases differs in color from that reflected from the surface; gold and copper, for example, are green when seen by transmitted light. 534. Metals can be obtained in a state of fine division and their color is determined by the size of the particles; metallic silver, for example, can be obtained by precipitation from solutions of its salts in forms which are red, green, yellow, etc. A convenient way to obtain metals in the finely divided condition is to produce sparks under water by alternately bringing into contact and sep- arating the terminals of an electric circuit. If the terminals are made of gold, each time they are separated and a spark is pro- duced, some of the metal is torn away in the form of particles so small that they stay suspended in the water and cannot be seen as such by the eye; in this condition gold is bright red in color and appears to be in solution. It is in what is known as the colloidal condition. When many substances insoluble in water are in an exceedingly finely divided condition they appear to dissolve and cannot be removed by filtration. Platinum in the form of a wire or sheet is a silver-white metal; in the colloidal condition it is black. Metals when in the finely divided condition exhibit prop- erties which are, no doubt, due to the fact that the surface of the metal is almost infinitely greater than when it is in the massive state. Iron, for example, reacts but slowly with the oxygen of 448 INORGANIC CHEMISTRY FOR COLLEGES the air when it is in the usual condition; it can be obtained so finely divided that it takes fire spontaneously when brought into the air. Finely divided metals serve as valuable catalytic agents; platinum is used to hasten the union of sulphur dioxide and oxygen in the manufacture of sulphuric acid, and nickel in the preparation of solid cooking fats by the addition of hydrogen to vegetable oils. 535. Change in State of the Metals. All the metals melt and can be vaporized, although in certain cases a temperature attain- able only in an electric furnace is necessary. The melting-points and boiling-points of the commoner metals are given in the table on page 443. The range in melting-points is great; mercury melts at 38.8, caesium, a metal that resembles sodium in properties, at 26.4 and tungsten at 3200. The use of mercury in thermom- eters and that of tungsten in electric lamps are based largely on the melting-points of the metals. Metals that boil below 1000 can be more or less readily purified by distillation. The fact that mercury boils at a much lower temperature than any other metal makes it possible to obtain it in a very pure condition for use in such measuring instruments as thermometers, barometers, and apparatus employed in the analysis of gases. Many metals dissolve in mercury and their presence even in small amounts makes it cling to glass and leave a film on its surface; it must, accordingly, be separated from these before it is used for the above purposes. The fact that mercury dissolves other metals and is readily volatilized is used in one process of extracting gold from its ores. The crushed ore is stirred with mercury and the solution of gold thus formed is then heated to the temperature at which mercury boils; after the latter has distilled off the residue is gold. When molten metals change from the liquid to the solid con- dition there is a change in volume; in most cases contraction takes place. Gold cannot be cast, because in passing from the liquid to the solid condition it contracts and, therefore, shrinks on solidi- fication from the mold into which it is poured. Since a sharp impression cannot be obtained in this way, in making coins, gold is struck with a die while hot. Bismuth expands on solidification. Antimony is put into lead which is to be used to cast type, because it forms with the metal an alloy that expands when it solidifies. 536. Vapor-density of the Metals. A few of the metals boil at a sufficiently low temperature to make it possible to determine PHYSICAL PROPERTIES OF METALS. ALLOYS 449 their vapor densities; in all cases that have been examined the results lead to the conclusion that the elements that are distinctly metallic in character, like sodium and mercury, are monatomic in the gaseous state. Antimony, which is more non-metallic than metallic is diatomic, and bismuth, which more nearly ap- proaches a metal in physical and chemical properties, yields a mixture of monatomic and diatomic molecules at a temperature just above its boiling-point, ALLOYS 637. The solids obtained when two or more metals are mixed in the molten condition and allowed to solidify are called alloys. On account of the fact that the presence of certain non-metallic elements, such as phosphorus and antimony, have a marked effect on the properties of the materials made in this way, these ele- ments are frequently added to the metals in the preparation of alloys. The so-called non-ferrous alloys only will be considered briefly here; those which contain iron, such as the various kinds of steel, are of such technical importance as to demand separate treatment when the chemistry of iron is discussed. The conditions in regard to solubility which exist when molten metals are mixed are the same as those obtained when substances liquid at the ordinary temperatures are brought together. It will be recalled that liquids either mix in all proportions, exhibit definite solubilities one in the other, or are practically immiscible. Exam- ples of these classes are, respectively, alcohol and water, ether and water, and kerosene and water. Similar relationships are found among the metals in the liquid state; examples of the three classes just mentioned are, respectively, silver and copper, lead and tin, and lead and copper. In the case when the metals exhibit a definite solubility, one in the other, the solubilities, as is the case in general with all types of solutions, vary with the temperature. All these relationships come into play when molten mixtures of metals cool and finally solidify, and they determine the physical structure of the solid alloy. In the case of the first class of alloys, namely, those made up of metals miscible in all proportions, the liquid mixture solidifies 450 INORGANIC CHEMISTRY FOR COLLEGES to a solid uniform in appearance when examined with a microscope of the highest power. The metals, apparently, are as evenly dis- tributed in the solid condition as they were in the liquid state; they are said to exist, accordingly, in solid solution. The melting- points of such alloys change gradually as the composition of the mixture is changed and are determined by the relative amounts of the metals present. Alloys of the third type are also simple in structure. In this case the immiscible mixture may be stirred to bring the constituents into more intimate contact and then allowed to solidify. On microscopic examination they show two distinct kinds of solid particles which are made up of the two metals present. When such an alloy is gradually heated to deter- mine its melting-point the lower melting metal first liquefies at its melting-point, and then the thermometer remains constant for a time as the heat supplied is used in melting the more fusible metal ; after this has taken place the thermometer rises until the melting- point of the second metal is reached, when it again comes to rest. Such alloys show, therefore, two distinct melting-points, which are those of its constituents. 538. The conditions that exist in the class of alloys made up of metals that show limited solubilities in each other are much more complicated and yield products more complex in structure. The alloys made from lead and tin, which are used as solder, are exam- ples of this class. Lead melts at 327, and tin at 232. When tin is added to molten lead, the temperature at which the mixture begins to solidify is lower than the melting-point of lead, and like- wise a mixture of tin and a small amount of lead begins to solidify at a temperature below the melting-point of tin. In both cases the melting-point is lowered a fact in accord with the generaliza- tion already given, namely, that the freezing-point of a solvent is lowered by the presence of a dissolved substance. In the accom- panying diagram (Fig. 32) are plotted the temperatures at which liquid mixtures of lead and tin begin to solidify. The significance of the diagram will be appreciated from the following discussion: The point marked x in the diagram repre- sents a mixture made up of 70 per cent of lead and 30 per cent of tin at the temperature 300; at this temperature the mixture is liquid. If the temperature falls, we see by following down the vertical line that at about 250 solidification begins; at this point PHYSICAL PROPERTIES OF METALS. ALLOYS 451 lead begins to separate in the solid, crystalline condition, 1 and this separation continues with falling temperature until 180 is reached, when the entire mass becomes solid. If we examine next the behavior of a mixture represented by the position y on the dia- gram one containing 80 per cent tin and 20 per cent lead at 250 we see that in this case tin 2 separates first at 190 and that the entire mixture is solid at 180. The mixture having the com- position indicated by the position of B on the diagram 31 per cent lead and 69 per cent tin solidifies completely at one tem- perature it has a definite melting-point. The mixture which has these properties is called the eutectic, the name being derived from the Greek words meaning " well-melting." f 6bO A 300 250 X .")(. Tin \ ^ "^~^Q. f^ y c 200 180 150 X -^^5 *-**^ 5 ^^ 180 Deg. D 1 ' ^* i E ) 10 20 30 40 50 60 70 &0 90 100 Per Cent Tin )0 90 80 70 60 50 40 30 EG 10 Per Cent Lead FIG. 32. The points in the area in the diagram above the curves AB and BC represent the composition and temperature of liquid mix- tures of the two metals; if these conditions place a mixture in the area between the curves AB and BD, it is composed of solid lead and the liquid mixture; and if in the area between CB and BE, of solid tin and the liquid mixture; below the line DE the alloy is a solid. It is evident from the above that an alloy made from a mixture of 20 parts of tin and 80 parts of lead will differ in properties from one containing 80 parts of tin and 20 of lead. The former will consist of crystals of lead embedded in the eutectic and the latter of crystals of tin in the eutectic. The physical structure of an crystals formed are a solid solution of about 4 parts of tin in 96 parts of lead. 2 The crystals contain 98 parts of tin and 2 parts of lead. 452 INORGANIC CHEMISTRY FOR COLLEGES alloy is rendered visible by polishing a flat surface of it and then treating the latter with a liquid which attacks it; if two or more substances are present they are affected differently, and after a short time the surface is found to be etched in such a way that the several ingredients of the alloy can be readily distinguished under a microscope. Figs. 33a, 336, and 33c are reproductions of photo- micrographs of alloys of tin and bismuth, which resemble in structure the alloys of lead and tin. Photographs of the former are used here to illustrate the microstructure of alloys because they can be reproduced more clearly than the photomicrographs of the lead-tin alloys. It is evident that such differences in physical structure will have a marked effect on the hardness, tensile strength, and other physical properties of the alloy. Returning to the consideration of the alloys of tin and lead, it will be seen from the diagram that if the one described, rich in lead (80 lead to 20 tin), is melted and allowed to cool, it begins to solidify at about 270. It will persist as a thick, viscous mixture made up of solid and liquid until the temperature falls to 180, when the mass solidifies completely; it will be in a plastic condi- tion over a range in temperature of 90 degrees, so that it can be worked or molded, as it has about the same viscosity as that of baker's dough. The alloy containing 80 parts of tin and 20 parts of lead stays liquid until a lower temperature is reached about 190 instead of 270, the temperature at which the alloy high in lead begins to freeze but is in the plastic condition during a drop in temperature of only 10 degrees, for the eutectic solidifies at 180. Such an alloy could be worked with at low temperatures, but if it is to be handled in the plastic condition, as in the case in soldering, it would solidify so rapidly that the manipulation would be very difficult. 539. The alloys of lead and tin are used as solders and we can now understand why mixtures of different compositions are used for different purposes. In soldering, two pieces of metal are joined by covering their ends with molten solder, and when the latter is in a plastic condition it is worked into shape with a hot tool, which is usually made of copper; a flux (521) is ordinarily used to keep the surfaces to be joined free from oxides so that they can alloy with the solder. The tensile strength of solder is greatest when it contains 72.5 per cent lead, but an alloy of this composition has a CO W o 28" 1 GO J> +J "8 fi i 1! I a] H fl 5 I PHYSICAL PROPERTIES OF METALS. ALLOYS 453 such a high freezing-point that it cannot be readily worked with a soldering tool. Alloys of lead and tin containing as high as 70 per cent of the former are used for coating sheet iron or steel for roofing, for castings, etc. Plumber's solder contains about 67 per cent of lead and 33 per cent of tin; it begins to assume the plastic state at about 245 and remains in this condition during a fall in temperature of 65 degrees. While in this condition the joint being made with the solder, between two pipes for example, is " wiped" that is, the plastic mass is molded into the desired form. 540. Ordinary soft or tinner's solder contains approximately 50 per cent of each metal. It can be worked at a low temperature (see diagram) and is used for making joints which do not have to be " wiped," such as those in tin cans. It is more expensive than plumber's solder because the price of tin is much higher than that of lead. In soldering tin, an alloy is used consisting of a mixture of 2 parts tin, 2 parts lead, and 1 part bismuth; the latter is added because it lowers the melting-point of the solder to 140 so that it can be used without risk of melting the tin. Plumber's solder begins to solidify about 40 degrees above the melting-point of tin and tinner's solder at only 10 degrees below this point. 541. The alloys of tin and lead have been discussed at some length because they are the simplest examples of alloys of this type that are used commercially. When three or more substances are used in making an alloy the conditions are much more complicated than those which have been described, and a wider range in proper- ties can be obtained. In the case of certain metals, compounds of definite composition are formed as the result of the union of the metals, and these act as distinct substances and by increasing the number of phases present add to the complexity of the resulting alloy. By selecting the right metals and Using them in the correct proportions, alloys can be made with any desired degree of hard- ness, ductility, toughness, tenacity, fusibility, etc. For many uses they are to be preferred to the pure metals. For example, sound castings can be made from copper with difficulty, and on account of the softness of the metal it cannot be filed and machined easily. Brass, on the other hand, which is an alloy of copper and zinc, can be molded and worked with facility. Small amounts of non-metallic elements are frequently added to metals that are to be used in making castings. When gases 454 INORGANIC CHEMISTRY FOR COLLEGES separate from molten metals on solidification they cause objection- able flaws in the casting. The separation of gases may be due to the fact that they were dissolved by the molten metal and lib- erated when it solidified, or that they were present in the impure metal as substances that interacted at the high temperature to form a gas. Commercial copper, for example, contains traces of copper sulphide and copper oxide, which interact at high temperatures to form copper and sulphur dioxide; when it solidifies the gas pro- duced in this way is expelled. If certain non-metals are added to the molten metal, they unite with the oxygen present and form oxides that are not decomposed at the temperatures used; phos- phorus and silicon are used for this purpose. Boron has recently been found to be of particular value in making sound castings from copper (518). An excess of the non-metallic element above that required for deoxidation generally results in increasing the hardness and brittleness of the alloys in which it is used. Active metals whose oxides are stable at high temperatures are also used as deoxidizers; among these are aluminium, magnesium, and manganese. The crystalline structure of alloys makes them susceptible to mechanical working and heat treatment; the changes effected by these agencies are more marked than those referred to in the case of metals (533). 542. The peculiar physical structure of alloys is utilized in making so-called anti-friction alloys which are used as bearings in machinery. These alloys consist of crystals of a hard constituent imbedded in a soft matrix; as a result, when they are used in the bearings for a steel shaft the hard material readjusts itself in the matrix to any inequalities in the shaft, and friction is thereby reduced; the soft material is worn away at the surface and the hard grains bear the pressure. Lead-antimony alloys are used for this purpose. The eutectic contains 13 per cent of antimony; alloys containing a larger amount of the latter consist of crystals of antimony, which is hard, embedded in the softer eutectic. Bab- bitt metal, which is a tin-antimony-copper alloy, is more expen- sive, but is to be preferred to the simpler alloy. The composition of a few alloys of technical importance is given in the table on page 455. The figures given in the table refer to the several alloys commonly used. Since there are changes in properties PHYSICAL PROPERTIES OF METALS. ALLOYS 455 with change in the composition of alloys, a great variety of sub- stances are used under the same name. For example, a large number of different brasses are made; type-metal varies widely in composition, etc. COMPOSITION OF SOME COMMON ALLOYS Alloy Tin Copper Zinc Lead Anti- mony Misc. Babbit metal 889 3 7 7.4 Bell metal 22 78 Brass 95-60 5-40 Brittania metal 45.5 1.5 40 13 Bronze for bearings 5 64 30 Ni, 1 Bronze, for castings 10 90 German silver . . 60 25 Ni, 15 Gold coins, U. S 10 Au, 90 Gun metal ... 9 91 M^onel metal 28 Ni, 69 M^osaic gold 65 35 Fe, 3 Nickel Coins, U. S 75 Ni, 25 Pewter 80 20 Rose's metal 22.9 27.1 Bi, 50 Silver Coins, U. S 10 Ag, 90 Silver sterling 2.5 Ag, 97 5 Solder, plumber's 33 67 Solder soft 50 50 Stereotype metal 10 2 70 18 Type metal 20 60 20 Wood's metal 14 24 Bi 50 Cd, 12 EXERCISES 1. Name several uses of a number of metals based on the fact that they are (a) hard, (6) soft, and (c) have a high heat conductivity. 2. (a) Calculate the weight per meter of a copper wire and an aluminium wire each 1 mm. in diameter. (6) Calculate the electrical resistance of 1 meter of each wire, (c) Calculate the diameter of a wire of aluminium 1 meter of which has the same resistance as 1 meter of a wire of copper 1 mm. in diameter, (d) Calculate the weight of the two wires described in c. (e) To what extent would the cost of aluminium and of copper be a factor in deciding which metal to use in the transmission of electricity? 456 INORGANIC CHEMISTRY FOR COLLEGES 3. Which of the following elements would be good and which poor con- ductors of electricity: (a) sodium, (6) calcium, (c) tellurium, (d) radium, (e) iodine. 4. (a) Why is it more agreeable to drink a hot liquid from a cup made of porcelain than from one made of aluminium? (6) Would you expect a liquid to cool more rapidly in a spoon made of tin or of silver? Why? (c) What metal would be best for making lightning rods? 5. (a) What properties of a metal determine whether it could be used as the tip on a soldering iron? (6) Is the specific heat of the metal a factor? The value of the specific heat of copper is 0.091 and of iron 0.119. From this point of view which metal is preferable? (c) W T hy is the tip of a solder- ing iron relatively large? (d) Why is a steel soldering iron used with a solder containing a high percentage of tin? (e) What determines what kind of solder should be used in soldering galvanized iron? A solder high in tin is used in this case. What conclusion can be drawn from this fact? 6. State why the wiped joint of solder on a lead pipe has a frosted appear- ance and the joint on a tin can has not. 7. To what class of alloys would you expect to belong the following pairs of metals: (a) silver and gold, (6) lead and zinc, (c) copper and antimony? CHAPTER XXXII THE CHEMICAL PROPERTIES OF METALS. METALLURGY 543. The metals resemble one another in their general chemical behavior with other substances, but they differ markedly in activity. They also show differences in valence, and since the valence of a metal is an important factor in determining the chemical properties of the compounds containing the metal, the comparative study of these elements must be made from this point of view. We have already seen, in a general way, that the relative activ- ity of metals is indicated by their position in the electromotive series (252) , and that their valence follows from their position in the periodic classification of the elements. These two generaliza- tions will be of the greatest value in correlating the many facts in regard to the chemistry of the metals and their compounds, and they will be constantly used for this purpose. 544. Behavior of Metals with Other Elements. All the metals unite with the halogens to form salts ; all form oxides when heated in the air or oxygen except those which fall in the electro- motive series below mercury; oxides of these inactive metals are known, however, but they must be prepared from salts. Sulphur unites with all the metals to form sulphides; at high temperatures most of the metals react with carbon and silicon. 545. Behavior of the Metals with Water. In their relative activity with water, the metals fall approximately into the same order as they do in the electromotive series. The alkali metals, which constitute the first family in the first group in the periodic classification, react rapidly with water at ordinary temperatures, and hydrogen and the hydroxides of the metals are formed. The metals of the alkaline earths those of family 1, group 2 also decompose water at room-temperature, but less rapidly. Magnesium and aluminium react with cold water very slowly, but both metals can be activated by treating them with a solution of a 457 458 INORGANIC CHEMISTRY FOR COLLEGES mercury salt; the more active metals set free mercury and when they become coated with a trace of the latter they decompose water; hydrogen is evolved and the metals are converted, into hydroxides. Such a combination of metals is called a metallic couple. Zinc, which is farther down in the electromotive series, does not attack water at ordinary temperatures, but when it is coupled with copper, by treating the metal with the dilute solution of a copper salt, it reacts slowly with water. Lead, which is in the electromotive series just above hydrogen, passes into solution to a very slight degree as lead hydroxide, but as soon as a trace of the latter is present the reaction stops. The metals below hydrogen in the electromotive series do not decompose water. The rate at which metals react with water is determined by their activity and by the temperature; the metals of the alkalies and alkaline earths react more or less rapidly with water at room- temperature; magnesium decomposes water slowly at 100, zinc at a higher temperature, and iron at red heat. 646. Behavior of the Metals with Acids. The behavior of the metals when brought into contact with acids varies widely; it is a subject of great importance on account of the fact that it deter- mines in many cases what metals can and what ones cannot be used for a specific purpose. Ordinary air contains carbon dioxide and water, which act as an acid, and wherever coal is burned sulphurous acid and sulphuric acid are found in appreciable quantities in the air. These acids attack many metals and their presence leads to the corrosion of such metals when they are exposed to the air. A number of facts in regard to the action of acids on metals have already been given ; these will be restated here in connection with a more systematic discussion of the subject. The action of acids on metals is determined by a number of factors, the most important of which are the following: the chemical activity of the metal, the oxidizing power of the acid, the presence of water which deter- mines whether or not the reaction is one involving ions, the con- centration of the solution of the acid, the chemical properties and solubilities of the resulting salt, and the temperature. 547. Action of Non-oxidizing Acids. Aqueous solutions of non-oxidizing acids react with the metals in the electromotive series down to hydrogen. Under these conditions the reactions THE CHEMICAL PROPERTIES OF METALS. METALLURGY 459 are ionic, and hydrogen and a salt of the metal are formed. Changes in concentration and temperature have the effect that might be expected from a knowledge of the influence of these factors on the ionization of the acid used. Oxidizing acids, if sufficiently dilute, behave in the same way, because their oxidizing power decreases with increased dilution as the result of the con- version of molecules the active oxidizing agent into ions. Many metals have two valencies; iron, for example, forms two chlorides, ferrous chloride, FeCk, and ferric chloride, FeCla, in which the valence of the element is 2 and 3, respectively. Since iron in the second compound is in a higher state of oxidation, and since nascent hydrogen is an active reducing agent, we would not expect ferric chloride to be formed in the presence of the hydrogen set free when iron dissolves in hydrochloric acid; and this is the case. In general, when a metal forms two classes of salts in which it shows different valencies, the salt containing the metal in the lower valence is formed when it dissolves in dilute acids. Whether or not a metal dissolves in an acid is determined not only by its activity, but by the behavior with water of the product formed. Aluminium, a very active metal, which stands high in the electromotive series, is scarcely attacked by dilute sulphuric, nitric, and most other acids. We shall see later that aluminium is a trivalent element, and on account of this fact its hydroxide is a very weak base; its salts are, therefore, hydrolyzed in water, and when the metal is treated with a dilute solution of an acid the salt first formed is probably decomposed by the water present to form an insoluble basic salt, which protects the surface of the metal from further action. Lead is slightly attacked by cold dilute solu- tions of hydrochloric acid or sulphuric acid, because the chloride and sulphate of the metal are difficultly soluble in water, and the small amount first formed protects the metal from further action. Lead chloride is soluble in hot water, and for this reason the metal dissolves in a hot solution of hydrochloric acid; lead sulphate is practically insoluble in hot water and the metal is not attacked by a hot dilute solution of sulphuric acid. Since the chlorides of all the metals in the electromotive series down to lead are sol- uble in water, all these metals dissolve readily in hydrochloric acid. They all dissolve in dilute sulphuric acid except aluminium, and in dilute nitric acid, except aluminium and chromium. 460 INORGANIC CHEMISTRY FOR COLLEGES 648. Action of Oxidizing Acids. Whether or not a metal dis- solves in an oxidizing acid depends on the activity of the metal and of the oxidizing acid, the concentration 'of the latter, and the properties of the products formed. Since oxidations are essen- tially molecular reactions, the smaller the concentration of water present, which ionizes the acid, the greater is the activity of the oxidizing agent. All the metals in the electromotive series down to and including silver are oxidized by concentrated nitric acid and by concentrated sulphuric acid. In certain cases the results obtained with nitric acid appear to be anomalous. When iron is placed in fuming nitric acid it does not dissolve, but is changed in such a way that it does not exhibit the properties of the element in its normal condition. After treatment with the acid iron will not dissolve in dilute acids, or precipitate copper from solutions of its salts; it is said to be passive. If the metal is scratched or struck, it assumes its usual condition. Chromium and cobalt can be rendered passive in the same way. It is probable that these metals are converted on the surface into oxides which contain a large percentage of oxygen and do not show basic properties and, therefore, are insoluble in acids; as they cover the metal they prevent it from interacting with solutions of salts. Much work has been devoted to the study of passive metals and a final conclusion as to the cause of the phenomenon has not been reached. Aluminium is practically unattacked by concentrated nitric acid. In dilute solutions nitric acid acts towards metals largely in its capacity as an acid; hydrogen and the nitrate containing the metal in its lower valence are formed; iron and dilute nitric acid yield ferrous nitrate, Fe(NOs)2, and hydrogen. With increasing concentration the oxidizing property of nitric acid shows itself; with strong nitric acid iron forms ferric nitrate, Fe(NO3)3, and as the result of its acting as an oxidizing agent the acid is reduced to nitric oxide or other oxides of nitrogen. Concentrated nitric acid has little effect on cast iron, and its vapor less; as a conse- quence, retorts made of cast iron are used in manufacturing the acid from sodium nitrate, and care is taken to have the tempera- ture of the retort above the boiling-point of the acid. Since there is a slight action between the acid and the metal, the con- densers used are made of aluminium, glass, or other resisting non- metallic material. THE CHEMICAL PROPERTIES OF METALS. METALLURGY 461 The action of nitric acid on tin illustrates another point. The metal forms compounds in which it functions in its metallic char- acter and shows the valence 2; it also forms compounds in which the valence is 4, and in which it plays the part of an acid-forming element. Stannic oxide, SnO2, is an acid anhydride and yields salts known as stannates. Tin dissolves in dilute nitric acid and stannous nitrate is formed ; when treated with concentrated nitric acid the metal is oxidized to the dioxide, which does not dissolve in the acid, and as a result a white insoluble powder, SnO2,H2O is formed. Since the normal nitrates are all soluble and those of the weakly metallic elements are not hydrolyzed in the presence of a large excess of nitric acid, all the metals except those noted above dis- solve in the strong acid. Hot concentrated sulphuric acid oxidizes all the metals down to and including silver. The solubility of the sulphate formed has a marked effect on the extent to which the reaction takes place. Lead sulphate, for example, is practically insoluble in sulphuric acid containing water until the concentration of the acid reaches about 77 per cent; it is, accordingly, not appreciably attacked by acid up to this concentration because the sulphate first formed serves as a coating to protect the metal; it dissolves when heated with stronger acid because the sulphate formed passes into solu- tion. The use of lead in the construction of chemical apparatus in which sulphuric acid is to be used is limited by these facts. The behavior of iron with sulphuric acid is quite different from that of lead; the sulphate of the metal is soluble in water and, consequently, the metal is attacked by mixtures of the acid and water until the concentration of the former reaches about 77 per cent, when action practically ceases. Application is made of this fact in the concentration of the sulphuric acid made in the cham- ber process; the acid is concentrated in lead pans until it reaches 77 per cent and is then transferred to iron stills in which the con- centration to 98 per cent is continued; it has been found that cast iron serves best. 649. The Effect of the Presence of Oxygen on the Solubility of Metals in Acids. In the presence of air solutions of acids dis- solve certain metals that are not attacked by these acids in the absence of oxygen. It is probable that solution results from the 462 INORGANIC CHEMISTRY FOR COLLEGES fact that the metal is first oxidized by the oxygen and the oxide formed dissolves in the acid present. Copper, for example, which is below hydrogen in the electromotive series, does not react with dilute acids, but if oxygen is present it is more or less rapidly converted into a salt. Advantage of this fact is taken in one of the commercial methods for the preparation of copper sulphate. Scrap copper is allowed to stand with dilute sulphuric acid in contact with the air until it passes into solution. Although the reaction takes place slowly, it is more economical and convenient than dissolving the metal in hot concentrated sulphuric acid with the attendant loss of sulphur dioxide, the difficulty of handling the hot acid, and the necessity of using a non-metallic container for the latter. Copper reacts with dilute acetic acid in the presence of air; the basic salt formed, which is insoluble in water, is used as a green pigment under the name verdigris. Tin, which is low in the electromotive series, near hydrogen, is scarcely attacked by dilute solutions of weak acids; in the presence of oxygen, however, it is dissolved. 550. The Corrosion of Metals in the Presence of Air. The metals down to and including copper are more or less affected by contact with the air; either an oxide is formed or, in certain cases, the presence in the air of carbon dioxide leads to the production of insoluble basic carbonates. Whether or not the metal is appre- ciably corroded is determined by the physical properties of the rust formed. In the case of aluminium, for example, the super- ficial coating of hydrated oxide adheres firmly, and as it is very thin it scarcely alters the appearance of the metal; with mag- nesium, however, another active metal, the result is different, for the basic carbonate formed is light and does not adhere so firmly, and, as a consequence, the metal is slowly disintegrated. Nickel, cobalt, and tin are affected to only a slight degree if at all; zinc becomes covered with a thin coating of basic carbonate, which resists further action and is, therefore, permanent in the air. Lead becomes coated with a very thin adhering layer of basic carbonate and resists further corrosion. A clean surface of copper oxidizes rapidly in the air, but the coating of the oxide formed is so thin it only slightly deepens the color of the metal; it is very permanent in ordinary air. THE CHEMICAL PROPERTIES OF METALS. METALLURGY 463 The behavior of iron is quite different; in this case the oxide formed is a powder the density of which is less than that of iron; when it is formed it occupies more space than the metal from which it was produced, and, therefore, falls away from the sur- face, leaving freshly exposed metal to be acted on farther; and the rusting continues more or less rapidly. A great deal of attention has been devoted to the study of the corrosion of iron on account of its technical importance; some of the more important conclusions reached will be discussed later (755). In the presence of the air of cities, especially near places where large quantities of coal are burned, the corrosion of metals takes place much more rapidly than in the pure air of the country. The sulphuric acid in the air, which is produced from the sulphur in coal, attacks the metals in the presence of oxygen, and salts are formed which are eventually converted in the presence of water and carbon dioxide into basic carbonates. The green color of bronze statuary kept out of doors and of copper roofing is due to this cause. In the neighborhood of the sea, copper and even nickel, which resists ordinary air, become covered with a coating of green basic car- bonates. The corrosion under these circumstances is due to the fact that the ocean and the air in its neighborhood contain salts that hydrolyze slightly and furnish a trace of hydrochloric acid which attacks the metal. Hydrogen sulphide is present in the air under certain condi- tions, and as it is decomposed by most metals with the formation of hydrogen and a metallic sulphide, its presence leads to the tar- nishing of bright metallic surfaces. Copper and silver act rapidly with the gas, and as the sulphides of these metals are black they are quickly tarnished. When the layer of sulphide formed on silver is very thin it gives to the metal a yellow color. 551. The facts enumerated above in regard to the corrosion of metals are taken into account in the use of metals for indus- trial purposes. Copper and zinc are used in the construction of buildings when it is necessary to expose a metal to the air. Roofs were formerly covered with copper or lead, but the high price of the metals has resulted in the substitution of cheaper material for this purpose. The question of cost has been met in many cases by covering iron or steel with a coating of a resistant metal. Gal- 464 INORGANIC CHEMISTRY FOR COLLEGES vanized iron is made by dipping iron into molten zinc ; the metal that adheres firmly to the iron if it has been carefully cleaned, serves to protect the latter for a long time from corrosion. In the process known as " sherardizing " objects made of iron, after treatment with dilute sulphuric acid to remove all oxide, are baked in zinc dust, and, as a result, are covered with a layer of the metal. Tin-plate is prepared by dipping sheet-iron into molten tin; iron is also coated with lead in a similar way. 552. Nickel, silver, zinc, gold, and copper are applied to the surfaces of other metals by the process of electroplating, which will be described later. Other substances than metals are used as protective coatings to prevent corrosion. Transparent lacquers and varnishes made from organic gums are frequently used to protect iron and silver. When iron is heated to a high tempera- ture with steam it reacts and an oxide of the formula Fes 04 is formed (736). This fact is used in covering articles made of iron or steel with the oxide, which acts as a protective coating; when they are heated in a closed retort and subjected to the action of high-pressure steam there is formed a surface layer of the oxide, which has a blue color. The process is used with rifle barrels, pistols, the so-called Russia or black sheet iron, etc. Paints of various kinds are also used to protect iron from corrosion. The protection against corrosion afforded by protective coatings is dependent on the physical nature of the latter; if they are porous or contain minute holes, air will reach the metal underneath and corrosion will take place more or less rapidly. It requires skill and attention to details to galvanize or tin iron properly and, as a consequence, many commercial articles of iron protected in this way soon rust. 553. As the result of the intensive study of alloys in recent years many substances of this class have been made that resist corrosion by the air and chemicals in general, and they are exten- sively used to replace metals. An alloy of iron and silicon known as duriron is particularly resistant and has found many uses, especially in chemical factories. The art of enameling metals has also developed, recently, and many large pieces of apparatus constructed of metal are lined with this type of material, which resists the action of solutions of most chemicals. Of the chemicals ordinarily used, hydrochloric acid is, perhaps, the most difficult THE CHEMICAL PROPERTIES OF METALS. METALLURGY 465 to protect against, because it is a very active acid and forms sol- uble salts with most metals. 554. The Action of Metals on Alkalies. The hydroxides of the metallic elements are bases, for they dissolve in acids and form salts; but those of certain metals act also as acids, dissolve in bases, and form salts in which the metal plays the part of an acid-forming element. For example, zinc hydroxide can function either as an acid or a base; with hydrochloric acid it forms zinc chloride and with sodium hydroxide, sodium zincate, Zn(ONa)2 or Na 2 ZnO 2 . Most metals of this class react with active bases; examples of these are aluminium, zinc, tin, and lead. Arsenic and antimony, which are essentially acid-forming elements, are also soluble in alkalies. Certain metals which do not react with solutions of bases do react with the alkalies at the temperature of fusion; platinum, for example, is converted slowly under these conditions into a platinate, and gold into an aurate. Other metals that resist fused alkalies are dissolved if an oxidizing agent, such as sodium nitrate, is present. " Owing to the presence of the latter the metal is oxidized to its higher valence in which it possesses acidic proper- ties, and it passes into solution as a result; manganese, chromium, and iron, examples of metals that behave in this way, yield man- ganates, chromates, and ferrates. Nickel does not show acidic properties in any of its compounds and as it melts at a relatively high temperature it is used in making vessels to be employed in the preparation of compounds involving fusion with an alkali and an oxidizing agent. If there is no chance for oxidation, iron can be used for the construction of apparatus in which fusions with alkalies are carried out. 555. Action of Metals on Salts. The replacement of one ele- ment by another when a metal is placed in a solution of a salt has already been discussed (252). It will be recalled that whether or not a reaction takes place is determined by the relative tenden- cies of the two metals to form ions. In general, when there are no disturbing factors present, a metal will displace a second metal from its salts if the first metal is above the second in the electromotive series. A number of applications are made of this property. For example, pins, which are made of brass or iron, are coated with tin by placing them in a solution of a tin salt. The 466 INORGANIC CHEMISTRY FOR COLLEGES action of metals on fused salts is utilized in isolating certain metals that cannot be readily separated from their oxides by reduction with carbon. Aluminium was first prepared in this way by heating sodium with fused aluminium chloride. When salts are in the molten state they show the electrical properties that they exhibit in aqueous solutions; they conduct electricity and enter into reactions similar to those which take place as the result of the presence of ions. As a consequence, the replacement of metals, one by the other, when metals are brought into contact with fused salts, resembles closely the replacement when solutions of salts are used. 556. Occurrence of the Metals. F. W. Clarke of the U. S. Geological Survey has calculated from all available data the probable composition of the earth's crust and atmosphere as far as it has been possible to examine it on its surface, at great heights in the air, and at the depth obtained in deep mines. The results are given in the following table: COMPOSITION OF THE EARTH'S CRUST Per Cent Per Cent Oxygen 49 . 85 Titanium 0.41 Silicon 26 . 03 Chlorine . 21 Aluminium 7 . 28 Carbon . 19 Iron 4. 12 Phosphorus 0. 10 Calcium 3. 18 Fluorine 0. 10 Sodium 2.33 Sulphur 0.10 Potassium 2.33 Barium 0.09 Magnesium 2. 11 Manganese 0.08 Hydrogen 0.97 Nitrogen 0.03 . 99.51 From the above it is seen that the total of all the elements other than those listed is less than one-half per cent. Aluminium is widely distributed and occurs chiefly as aluminium silicate in association with other silicates, which make up the more important rocks. As the result of the disintegration of the latter under atmospheric influences, the soil is formed. As a consequence, it is clear why oxygen, silicon, and aluminium occupy their positions in the above table. The other metals listed in the table occur as silicates in igneous THE CHEMICAL PROPERTIES OF METALS. METALLURGY 467 rocks, and also in other forms, which owe their origin and existence to the chemical properties of the metals, and to the physical prop- erties of the salts formed from them. When the earth was being formed as the result of the cooling of a mass made up of elements in the form of vapor, the more active elements probably united first and the chlorides of the alkali metals and alkaline earths were soon formed. Since these compounds are soluble in water they have accumulated in the ocean and in salt beds formed as the result of the drying up of inland seas. Other soluble salts of the alkali metals occur in nature in arid regions; the sodium nitrate of Chile is, in all probability, the result of the nitrification of organic matter which was accumulated in one locality. Sodium car- bonate is found in certain alkaline lakes, and borax in desert regions. The alkaline earths occur as carbonates and sulphates, calcium carbonate being an important constituent of many rocks, such as limestone and marble; gypsum, CaSO4,2H2O, is also an important mineral. Magnesium occurs extensively as silicates, as the sulphate in salt deposits, and as the carbonate in rocks. As we pass to the less active metals in the electromotive series, in addition to silicates, we find carbonates, oxides, hydrated oxides, and sulphides as the chief minerals containing these ele- ments. When we reach copper in descending in the series, we find that the metal occurs in the above forms but is also found in the free condition in nature. Metallic iron, cobalt, and nickel are present in some meteors, but they do not occur in this form in the earth's crust. The elements below copper in the series all occur in the native state, that is, not in combination with another element. The more active of these are also found in combina- tion, sulphides being the most abundant of these compounds. Silver chloride is an important mineral. The noble metals occur almost exclusively in the metallic condition; gold, however, occurs as the sulphide and telluride. METALLURGY 557. The science which treats of the methods used to obtain the metals in the free condition from the compounds that occur in nature, is called metallurgy. All the naturally occurring com- 468 INORGANIC CHEMISTRY FOR COLLEGES pounds cannot be used as sources of the metals on account of the difficulties encountered in certain cases. Potassium, for example, is widely distributed in the form of feldspar and other silicates, but these cannot be used economically as a source of the metal or its compounds; potassium chloride, however, can be readily used for these purposes. The minerals, in more or less pure form, which are actually used as sources of the metals, are known as ores. These are commonly oxides, sulphides, carbonates, and chlorides compounds simple in composition, from which the metal can be extracted by the use of a few chemical reactions comparatively easy to carry out. The process employed in the isolation of any particular metal is determined by the chemical properties of the compound used and those of the impurities present in the ore. If we take iron as an example, its chief ores are oxides of the metal mixed with sand and other silicious material. Since the oxides of iron are reduced to the metal by carbon at a high temperature, the ore is mixed with coke and heated in a blast furnace to the tem- perature at which iron melts. A blast of air is blown into the furnace during the reduction to burn a part of the coke and thus produce the heat required. Since, at the temperature used, the silicon dioxide present in the ore does not melt and in the solid form prevents contact between the oxide of iron and the reducing agent, calcium carbonate is added to the charge of ore and coke; it forms a silicate with the silicon dioxide, and all the silicious material melts. The calcium carbonate is, thus, used as a flux; the molten material other than the iron is called the slag. A number of metals which occur as oxides are obtained by reducing their ores with carbon; among these are zinc, iron, tin, and bismuth. 558. When the ore to be used is a sulphide, it is converted into an oxide before reduction; carbon does not reduce sulphides, because the reaction between carbon and sulphur is endothermic. The ore is first heated in the air, or roasted, as a result of which it is changed into an oxide, and sulphur dioxide is given off; the oxide is then reduced by carbon. Zinc is obtained in this way from its sulphide, sphalerite, ZnS. Other elements obtained in this way from sulphides are cadmium, cobalt, and nickel. THE CHEMICAL PROPERTIES OF METALS. METALLURGY 469 In the case of certain metals whose oxides are readily reducible, it is not necessary to use carbon as a reducing agent when the ore is a sulphide. Advantage is taken of the fact that the oxide is reduced by the sulphur in combination with the metal as sulphide. For example, when copper sulphide is heated with copper oxide the two compounds interact and copper is formed: CuS + 2Cu 2 O = 5Cu 4- SO 2 In the case of metals of this kind copper and lead are examples the ore is roasted until only the proper amount of itis converted into oxide; it is then heated to a higher temperature and as the result of the interaction of the oxide and sulphide the metal is obtained in the free condition. 559. Silver usually occurs with lead or copper; its separation from these metals after extraction from their ores will be described later (713). The isolation of mercury from its sulphides is easy to effect because they are converted into the oxide when roasted, and the latter readily breaks down at a higher temperature into the metal and oxygen. All that is necessary is to heat the ore in air and condense the vapor of the metal formed. The metals of the platinum group are found free in nature more or less alloyed with one another; they are separated from one another by converting them into salts which differ in solubility. Gold occurs native, in the sulphide ores of lead and copper, and as a telluride along with silver. Since its metallurgy consists largely in its separation from other metals it will be considered later. 560. The metals which have been discussed up to this point are those whose oxides are reducible by carbon at temperatures readily obtained when the latter burns. We have seen that the methods used in the isolation of these metals varied with the stability of the oxides, and that they changed, in general, progress- ively with the position of the metal in the electromotive series of the elements. In the case of metals whose oxides are not reducible by carbon at the temperature produced when the latter burns and, therefore, obtained in furnaces (about 1800), other methods have to be employed. Either the reduction is carried out in an electric arc furnace, where the temperature reaches 470 INORGANIC CHEMISTRY FOR COLLEGES about 3500, or the metal is isolated by the electrolysis of one of its compounds in the molten condition. In certain cases a fused salt of the element is heated with a more active element, which has been previously prepared by electrolysis. Calcium can be obtained by reducing its oxide with carbon in an electric furnace, but as the metal unites with carbon to form calcium carbide, CaC2, care has to be taken to avoid an excess of the reducing agent. The metal is ordinarily obtained as the result of the electrolysis of fused calcium chloride; magnesium is isolated in a similar way from a fused double salt, MgCl2,KCl. Double salts are frequently used in the electrolysis because they melt at a lower temperature than the simple compounds. Sodium is obtained by electrolyzing fused sodium hydroxide; and potas- sium from the fused chloride. Aluminium, which is of such technical importance, is prepared on the large scale by the electrolysis of a solution of its oxide in fused cryolite, AlFs,3NaF. Further details of some of these electrolytic processes will be given later, EXERCISES 1. Tabulate the properties of the common metals given in Chapter XXXII in the following way: Prepare a table in which the following metals arranged according to their position in the electromotive series form the first vertical column: Na, Ca, Mg, Al, Zn, Fe, Sn, Pb, Cu, Hg, Ag, Au. Place horizontally the following headings: (a) heated in air, (6) air at room temperature, (c) water, (d) solutions of non-oxidizing acids, (e) oxidizing acids, (/) solu- tion of alkalies, (gr) occurrence in nature, (h) metallurgy. Draw lines hori- zontally between the syrrbols of the elements and vertically between the headings and fill in each square formed with a brief statement which gives the facts called for. 2. Making use of only the facts given in this chapter, how could you dis- tinguish the following from each other: (a) Mg and Ag, (6) Cu and Au, (c) SnandZn, (d) CaandPb, (e) Al and Zn, (/) Al and Sn, (g) PbandZn? CHAPTER XXXIII ELECTROCHEMISTRY 561. Electrical energy is used extensively in the preparation of a large number of compounds which are of great industrial im- portance; and its use is being constantly extended as cheap sources of electrical power are made available and the price of coal rises. The up-keep and utilization of an electrical installation using water-power is small compared with one requiring fuel as a source of energy, for after the first cost has been met, the energy con- sumed is derived from falling water. Electrical energy is used in chemical operations as a source of heat in furnaces of various types, and is directly converted into chemical energy in the prep- aration of metals from their ores, and in the manufacture of many important compounds. This branch of industry has become so important and the knowledge required to conduct and develop it so detailed that it has led to a new type of specialist known as an electrochemical engineer. The conversion of electrical energy into chemical energy can be carried out quantitatively in certain cases in either direction, and as we cannot measure chemical energy directly, it is necessary to use such transformations in order to get an accurate knowledge of the energy changes in chemical reactions. We have seen how important is a knowledge of these changes in energy, and a chemist must, therefore, be familiar with those parts of the science of electricity which are used in interpreting chemical change. From the practical point of view, and from the theoretical aspect of the subject, which is the basis for applications of elec- tricity in chemical industry, a knowledge of electrochemistry is a necessity. The nature of the electric current and the methods used to measure it will be first briefly considered, and then its simpler applications in industry and the pure science will be dis- cussed in some detail. 471 472 INORGANIC CHEMISTRY FOR COLLEGES 562. The Nature of the Electric Current. In the modern con- ception of the constitution of matter which has been referred to briefly (398), the view is held that the atom is made up of a posi- tive nucleus surrounded by negative charges of electricity, called electrons. Under certain influences these electrons can pass pro- gressively from one atom to another, and when their motion takes place in a single direction the phenomenon known as an electric current is observed. According to this view, the effect produced by such a current is due to a stream of negative charges of elec- tricity moving in a single direction. Since different substances differ markedly in electrical conductivity, it is concluded from the above hypothesis that they differ in the ease with which their atoms give up electrons. It will be recalled -that according to the electronic conception of valence (442), when chemical union takes place between two elements, an electron passes from the positive or metallic element to the negative or non-metallic element the metallic element loses the electron. This view is in accord with the fact that metals conduct the electric current, if the hypothesis in regard to the nature of the current is accepted. The electrons present in metallic atoms can, with greater or less ease, be set in motion and move progressively through the metal under certain influences. Before these are discussed, however, it is advisable to consider more fully electrical energy and how it is measured. 563. Measurement of Electrical Energy. This type of energy, like all others, has two factors, intensity and quantity. A rough analogy exists between the flow of water through a pipe and the flow of an electric current through a wire. The mechanical energy of the moving water is determined by the amount of water in motion and the pressure on it which makes it move. If a turbine is driven by the water, the amount of energy which can be trans- formed into work depends on these factors; it is for this reason that in the development of water-power a source is sought which yields a large amount of water the quantity factor where the fall is as great as possible, in order that a high pressure can be obtained the intensity factor. In an analogous way, when an electric current flows through a wire, the energy is determined by the quantity that passes and its pressure, which is called in this case electromotive force or poten- tial. The quantity factor of electrical energy is expressed in a ELECTROCHEMISTRY 473 unit called a coulomb 1 and the unit of the intensity factor is the volt; the electrical energy is the product of these two, and is expressed in joules. When electrical energy is being produced or used and the time factor in the transformation is considered, additional units, which express the rate at which the change takes place, are found to be convenient. The rate in the case of the quantity factor (coulombs) is measured in amperes; 1 ampere is 1 coulomb per second. The rate of transformation of electrical energy (joules) is expressed in watts; 1 watt is 1 joule per second; 1 kilowatt is 1000 watts; 1 kilowatt-hour is 1 kilowatt for 1 hour = 1000 X 60 X 60 joules. The use of these terms is comparatively common; the amount of electrical energy required to run a lamp is usually expressed in watts per candle-power, and the cost of electricity is expressed as a certain number of cents per kilowatt. Electric lamps are also marked with the voltage of the current which must be used in connection with them, for, as we shall see later, the heat developed when a current passes through a lamp and, con- sequently, the temperature of the filament and the light given off by it varies with the electromotive force of the current. In commercial work mechanical energy, such as that of a steam engine, is expressed in horse-power; 1 horse-power equals 746 watts. Instruments have been devised which can be used to measure amperes and volts; it is only necessary to include in an electric circuit an ammeter and a voltmeter to be able to read off on the scale of the instrument the values of the quantities involved. Instruments are also used which measure coulombs and watts. For the values of these quantities expressed in absolute units the reader is referred to books on physics; but an appreciation of their significance can be obtained by noting the values in the case of certain common electrical appliances. The current commer- cially furnished for lighting and power has an electromotive force of 110 or 220 volts. Electric arc lights are designed to take from 10 to 250 amperes according to the brilliancy of the light required. A tungsten incandescent light consumes about 1.2 watts per candle-power. An ordinary dry cell has an electromotive force 1 Coulomb (1736-1806), Volta (1745-1827), Ampere (1775-1836), Ohm (1787-1854), Watt (1736-1819), and Joule (1818-1889), were distinguished physicists who contributed much to our knowledge of energy and electricity. 474 INORGANIC CHEMISTRY FOR COLLEGES of about 1.6 volts, and a single cell in a storage battery about 2 volts. 564. Methods of Producing a Current of Electricity. Since an electric current possesses energy, it is necessary to transform some other kind of energy in order to produce it; mechanical energy, heat, or chemical energy can be used for this purpose. From Mechanical Energy. If two substances are brought into contact and then separated, they will be found to be charged with electricity, provided they are insulated so that the charges pro- duced cannot be lost by conduction. The charge produced by a single contact is excessively small, but if the objects are brought together many times, and one of them is discharged after each contact, the charge developed on the other can be made appre- ciable. So-called frictional electricity is produced in this way. The usual method of converting mechanical energy into elec- trical energy is by means of a dynamo. When a magnet is passed through a circle made of metallic wire, a current of electricity flows through the wire as long as the magnet is in motion; and the reverse of this is true also; when an electric current is passed through a coil of wire surrounding a magnet the latter is set in motion. These fundamental facts, discovered by Faraday, were utilized in the invention of the dynamo, which converts the mechan- ical energy of moving magnets into an electric current, and of the electric motor, in which magnets are made to move as the result of the passage of an electric current around them. 565. From Heat Energy. When two metals are brought into contact, they each become charged with electricity, as we have seen, but no current flows the electricity is static. An electric current can be set up, however, when two metals are in contact, if heat energy is supplied. This can be done in a simple way. If two wires of different metals are connected at their two ends and one of the points of contact is heated, a current of electricity flows through the wires as long as the two junctions are at different temperatures. The heat supplied furnishes the energy necessary to separate the electrons from the metals and cause them to move in the wires. The amount of heat energy changed into elec- trical energy increases with rise in temperature. This is the result, according to the accepted hypothesis, of an increase in the number ELECTROCHEMISTRY 475 of electrons set in motion and the greater intensity with which they are given off. The proportion of the heat changed into electricity in this way is very small, and the method is not used as a means of making electrical energy. It is utilized, however, in an important instru- ment for measuring high temperatures, known as the thermoelectric pyrometer, which is much used in chemical work. The temperatures to which the instrument is to be subjected determines the metals used in the pyrometer; if these are high, one wire is platinum and the other an alloy of platinum and rhodium or iridium, which are rare metals of the platinum group. Instruments using these metals are durable but expensive, and since the metals are so alike chemically the electromotive force of the current set up when they are heated is very small. If lower temperatures are to be measured, metals farther apart in the electromotive series can be used and a greater difference of potential is obtained. Combinations com- monly used are platinum and silver, and iron and chromel, which is an alloy of nickel and chromium. In constructing a pyrometer, the ends of the two wires used are fused together and the junction protected by a tube of fused silica or porcelain, one end of which is closed. This end is placed where the temperature is to be measured. The other ends of the wires are kept at a constant temperature in a vessel packed with cotton wool, or, when extreme accuracy is desired, in melting ice. To these ends is attached a millivoltmeter an instrument designed to read thousandths of a volt. The pyrometer is first calibrated by placing its terminal into sub- stances at known temperatures, such as boiling sulphur and certain metals at their melting-points. The reading on the voltmeter is noted in each case, and the results plotted, using tem- peratures and millivolts as co-ordinates. From the curve drawn through these points the temperature at any reading on the volt- meter can be read off. Pyrometers of this kind are of the greatest value in the determination of high temperatures, and as many industrial operations are carried out at these temperatures the instrument is much used. The thermoelectric pyrometer was invented by Le Chatelier, who first discovered the law of mobile equilibrium; it is frequently designated by his name. 476 INORGANIC CHEMISTRY FOR COLLEGES 566. From Chemical Energy. If a rod of zinc and a rod of copper are placed in a dilute solution of hydrochloric acid, the metals not being in contact, we observe that the zinc reacts with the acid and hydrogen is given off at the surface of the metal; the copper is not affected. If the temperature of the solution is carefully observed it will be found that it rises as the reaction proceeds chemical energy is being transformed into heat. If, next, the two rods are connected outside the solution by a metallic wire, it will be observed that the hydrogen is now evolved from the copper rod and not the zinc; and if a proper instrument for detecting an electric current is included in the wire circuit, it will be found that a current is flowing in it. The temperature change in this case is very small. The chemical products of the action in both cases are the same; zinc passes into solution as zinc chloride, hydrogen is set free, and the copper is not altered. In the first case all the chemical energy lost when the metal passed into solution was converted into heat; in the second case a part of the energy was transformed into elec- tricity. The explanation offered of these striking facts is as follows: When a rod of zinc is placed in water there is a tendency for some of the atoms of the metal to pass into solution as positive ions, which, it will be recalled, are atoms which have lost one or more negative charges as electrons, and are, therefore, said to be positive. When this takes place the rod of zinc necessarily assumes a nega- tive charge because it holds the electrons which were formerly associated with zinc atoms which passed into solution. There is thus set up a difference of potential between the solution and the metal; the change soon comes to equilibrium when the zinc ions reach a certain concentration. The case is different if an acid is present. The metallic zinc is in contact with hydrogen ions hydrogen atoms each of which has lost one electron. Since zinc has a greater tendency than hydrogen to become an ion that is, lose negative electrons it gives up to the hydrogen ions the excess of electrons which have accumulated on it as the result of the passage of zinc as ions into water. When the hydrogen ions regain electrons they pass into atoms, form molecules, and hydrogen gas is given off. We see, as a result of the consideration of the reaction between ELECTROCHEMISTRY 477 a metal and an acid from the point of view of the conception of electrons, that whether or not a given metal dissolves in an acid is determined by the relative ease with which it and hydro- gen lose these negative charges of electricity. The transfer of the electrons from the zinc to the hydrogen atoms takes place with a liberation of energy, which appears as heat. As the electrons left on the metallic zinc when the atoms of the metal pass into solution are immediately transferred to the hydrogen ions, the metal does not assume an electric charge. The explanation can be continued to include the case where a rod of zinc in metallic contact with one of copper is placed in hydro- chloric acid. When pieces of copper and zinc are brought together in the absence of a solution, the zinc loses electrons to the copper and a difference in potential is set up. We have seen that a current can be produced if the metals are brought into contact at two places to make a circuit and one of the junctions is heated; zinc, the more positive element, loses electrons to copper. When a rod of copper is joined to one of zinc and the free ends are placed into a solution of an acid, the electrons left on the zinc when its atoms pass into solution, immediately flow through the wire to the copper. The excess of negative electrons is now on this metal, and it is from the copper that the hydrogen ions get the charges necessary to change them into atoms; the gas is, accordingly, liberated under these conditions from the copper electrode. The flow of negative electrons through the wire from the zinc to the copper produces an electric current. The energy change asso- ciated with the conversion of metallic zinc into ions and of hydrogen ions into the gas is largely the change of chemical energy into electricity. It should be carefully noted at this point that the terms used in connection with the flow of an electric current are somewhat confusing. Before the advent of the electronic conception of the nature of the current, it was the common practice to assume that the electric current flowed in a circuit from positive to negative. We are now of the opinion that the current is produced as the result of the flow of negative electrons in the reverse direction. The two views are directly opposite, but the older method of expressing the direction of the flow of the current is still used with the understanding that it refers to what is called the positive cur- 478 INORGANIC CHEMISTRY FOR COLLEGES rent. In the electric cell just considered the electrons flow in the wire outside the liquid from zinc to copper and this is, consequently, the direction of flow of the negative current; but in common prac- tice the current is said to flow in the opposite direction. It will be easy to avoid confusion if it is remembered that by definition the current is said to flow always in the direction opposite to that of the moving electrons, and that when the nature of the current is not specified this so-called positive current is understood. Using this method of expressing the direction of the flow of the current in the cell described, the current flows outside the solution from the copper pole, which is said to be positive and is called the cathode, to the zinc pole, which is said to be negative (the anode) ; in the cell it flows from the zinc to the copper and the circuit is thus completed. It will be well to remember that in the solution the positive current always flows from the metal that is passing into solution as positive ions to the metal upon which the positive ions are discharged. The direction of the flow of the electrons and, consequently, the negative current is the opposite of this. 567. Metallic Couples. The explanation by means of the electronic hypothesis of the way in which an electric current is set up, can be applied to the case of metallic couples (545). It was pointed out that zinc does not react appreciably with water, but that when it is brought into contact with copper the zinc-copper couple hydrogen is evolved. We have just seen that when zinc is put into water some of the metal passes into solution as ions. The hydrogen liberated is deposited on the metal and the reaction soon stops. If it is in contact with copper, the gas is given off from this metal and the zinc left free to react further; the reac- tion is slow because the zinc hydroxide formed protects the metal to some extent. Zinc does not react appreciably with a solution of sodium hydroxide, but hydrogen is given off from the solution if the zinc-copper couple is used, for under these circumstances the zinc hydroxide is dissolved by the alkali as soon as formed. Perfectly pure zinc will not dissolve to any extent in the hydrochloric acid; the metal soon becomes covered with hydrogen on its surface and reaction ceases. If it is impure it contains substances which immediately set up a difference of potential with the metal and the gas is evolved from these impurities, and ELECTROCHEMISTRY 479 the zinc left free to act. It was for this reason that a small amount of a solution of copper sulphate was added to zinc and the acid in one of our earliest experiments (47). The metal reacted with the salt, copper was deposited in contact with the zinc, and a couple was formed. If we desire to prevent impure zinc from acting with a dilute acid, the metal is coated with mercury. As a result an alloy of uniform composition is formed on the surface of the metal and there is no opportunity to set up local currents as is the case with the impure metal alone. It is for this reason that zinc rods and plates used in the electric cells are " amalgamated " with mercury; the metal is not attacked by the materials in the solution when the cell is not working. The couple formed when the surface of tinned iron is abraded so that the iron is exposed to the air acts in a similar way; an electric current is set up in the presence of water, the more positive element iron is converted into oxide, and hydrogen is given off on the tin. In the case of galvanized iron it is the zinc which is corroded, as it is the more positive metal. 568. An Electric Cell. The combination of zinc, copper, and hydrochloric acid previously described could not be used conveni- ently as a source of electricity, because after a time the copper electrode would be covered with bubbles of hydrogen, the zinc would dissolve in the acid in the ordinary way, evolving hydrogen on its surface, and the chemical energy would be transformed into heat. It is evident that hydrogen must not be given off at the cathode, and that there should not be enough hydrogen ions present to interact with the zinc when the cell is not being used to produce electricity that is, when the circuit is open. These conditions are brought about in a number of ways. In the Daniell cell the anode is zinc surrounded by a solution of zinc sulphate placed in a cup made of baked clay which has not been glazed and is, there- fore, porous; this rests in a solution of copper sulphate in which the cathode of copper is placed. It is necessary to separate the two solutions, because if the copper sulphate came into contact with the zinc a reaction would take place; zinc would pass into solution and copper would be deposited in the former. The porous cup prevents the mixing of the solutions, but as these penetrate its walls it is possible for the current to flow. When the two poles are connected outside the cell, the zinc passes into solution in the 480 INORGANIC CHEMISTRY FOR COLLEGES way described above at length and metallic copper is deposited on the copper electrode. The so-called gravity cell (Fig. 34) is an ingenious improvement on the one just described, as it avoids the use of a porous cup, which adds to the resistance of the cell and, therefore, reduces the current obtainable from it. The -copper electrode and copper sulphate crystals are placed at the bottom of a glass jar which is nearly filled with water, and the zinc electrode is supported near the top of the jar under the latter. When the terminals are con- nected a current is set up and after some time the cell is running ZnSO 4 Solution Cu S0 4 Solution Copper Sulphate Crusta/..- Copper FIG. 34. normally. What occurs then is as follows: The presence of the crys- tals of copper sulphate keeps the solution around the copper electrode saturated with the salt, which fur- nishes the copper ions deposited on this electrode. The sulphate ions move toward the zinc electrode and form zinc sulphate with the metal, which passes into solution; there is, thus, an accumulation of this salt around the anode. The two salts in solution are kept from mixing for a long time by the fact that a saturated solution of copper sulphate has a greater density than one of zinc sulphate, and stays at the bottom of the cell if it is not disturbed; it is for this reason that it is called a gravity cell. In the so-called dry cell, which is so extensively used, zinc is the anode, ammonium chloride and zinc chloride in water the elec- trolyte, and powdered manganese dioxide and carbon in contact with a rod of carbon the cathode. When a current is set up in the cell the ammonium ions liberated at the cathode break down into ammonia and hydrogen. The latter does not accumulate on the cathode, but is oxidized to water by the manganese dioxide. The ammonia unites with the zinc salt in solution and the satura- tion of the latter with the free gas is thus avoided. The cell is constructed in such a way that it is very compact and efficient. The container is the zinc which serves as one electrode. This is covered on the inside with blotting-paper in order to separate the metal from the powdered manganese dioxide, which is packed ELECTROCHEMISTRY 481 in tight, after a carbon rod to serve as a contact has been placed in the center of the cell. A solution of ammonium chloride con- taining zinc chloride is next poured in to fill the interstices between the grains of the manganese dioxide, and the cell is finally closed with melted pitch. The cathode of the dry cell is manganese dioxide; it is in the form of a closely packed powder which nearly fills the entire cell and, accordingly, has a very large surface; it is separated from the zinc by the thickness of a sheet of blotting- paper and, as a result, the internal resistance of the cell is small. These conditions make it possible to get a large amount of current from such a cell in a given time a dry cell connected directly with an ammeter will develop as high as 30 amperes. The large surface of the cathode prevents the accumulation of hydrogen on it, and the cell can be used continuously. If a gas collects on an electrode it prevents to a greater or less extent contact between the solution and the electrode; under these conditions the elec- trode is said to be polarized. In the Leclanche cell the electrolyte is ammonium chloride, and the cathode is a mixture of manganese dioxide and carbon that has been pressed into the form of a hollow cylinder; the anode is a rod of zinc. The chemical reactions involved are the same as in the dry cell, but as the cathode presents a rela- tively small surface hydrogen soon accumulates on it and as polari- zation takes place the current soon drops off. The cell can be used only intermittently. In the cells which have been described, the electricity is generated as the result of utilizing the energy set free when a metal passes into solution as positive ions and the positive ions of another substance lose their charges and come out of solution in the free condition for example, Zn + Cu + >Zn + " + Cu. For this reason cells of this type are called replacement cells. Combina- tion cells based upon the direct combination of two elements, oxi- dation cells, and concentration cells in which the electrodes are the same metal in contact with different concentrations of one of its salts, have all been fully studied and are of interest; but their consideration in a very elementary account of electrochemistry is inadvisable. 569. The Quantity Factor of the Energy in Electrochemical Change. In the consideration of the electric cell up to this point 482 INORGANIC CHEMISTRY FOR COLLEGES attention has been paid to the way in which the current is pro- duced, but nothing has been said as to the quantity of electricity formed or its intensity factor. The relation between the amount of electricity produced and the weight of the metal involved is best determined by studying the conversion of metallic ions into free metals when a current of electricity is passed through solu- tions of their salts. This process is just the reverse of that by which a current is formed the metallic ion loses its positive charge as the result of the taking up of electrons and is thereby changed to the metallic condition. These changes can be indi- cated as follows where represents an electron a negative charge: Zn + + 2 ;= Zn. Read from left to right the equa- tion indicates that a zinc ion takes up two negative charges and is thus changed to the free metal. This is what happens when zinc is deposited by electricity from a solution of one of its salts. Read in the opposite direction the equation means that the metal loses two electrons and becomes an ion; the change indicated is the production of electricity. The action of the electric current on solutions was studied by Faraday in 1834, and important facts were discovered. It was shown, first, that the quantity of any one substance decomposed by the electric current is proportional to the quantity of elec- tricity passed through its solution, and, second, that the quanti- ties of two or more substances liberated by equal quantities of electricity are proportional to the chemical equivalents of these substances. This statement is known as Faraday's law. By chemical equivalent is meant the weight equivalent to 1.008 grams of hydrogen. This is 1 gram-atomic-weight of a univalent ele- ment, J gram-atomic-weight of a bivalent element, etc. This law is of fundamental importance and is the basis for all calcu- lations involving the relationship between quantities of elec- tricity, measured in coulombs, and weights of substances in grams. It was found by experiment that 96,500 coulombs are required to liberate 1 gram-atomic-weight. Accordingly, this quantity of electricity will liberate 1.008 grams of hydrogen, 35.46 of chlorine, 23.00 grams of sodium, 65.37 -^ 2 = 32.68 grams of zinc, 27.1 -T- 3 = 9.03 grams of aluminium, etc. Since reference is often made to the quantity of electricity which liberates one equivalent of an ELECTROCHEMISTRY 483 element, 96,500 coulombs, it has been given a special name and is called a faraday. Under carefully controlled conditions the deposition of metals by the electric current from solutions of their salts can be carried out with great accuracy, and this method is used in an instrument called a coulometer to measure quantities of electricity. It is simple in construction. Two pieces of copper placed in a solution of a copper salt are connected with the poles of the source of electricity to be measured. As the current passes through the instrument the metal is dissolved at one pole and deposited at the other. The increase in weight of the pole connected with the anode is a measure of the current passing through the instrument; for each gram- atomic-weight of copper deposited, 63.57 grams, 2 X 96,500 coulombs of electricity passed. If the time required for the depo- sition is noted the amperes can be calculated. We have just seen the relationship between the quantities of electricity and weights of substances which are involved in the change represented by the equation Zn + + + 2 = Zn. If the reac- tion is carried out in the reverse direction and the zinc is changed into ions under the conditions which permit the formation of an electric current, as in the gravity battery, the quantity of elec- tricity produced when a given weight of the metal dissolves, is just equal to that required to deposit the same weight of metal from its ions. Thus, 1 gram-atomic-weight of a univalent element in passing into the ionic condition will liberate 96,500 coulombs; likewise one-half of a gram-atomic-weight of a bivalent element will liberate this quantity, etc. In brief, each valence of an element when it appears or disappears involves a transfer of 96,500 coulombs, when we consider gram-atomic-weights in all cases. We are now in a position to understand the quantity of elec- tricity that can be obtained from a given cell. This is deter- mined by the weight of the element which is dissolved; each gram- equivalent furnishes 96,500 coulombs. The time-factor involved is an important one; two cells may furnish the same amount of electricity, but one may do this in a few minutes and the other require hours. The rate at which the current is furnished is deter- mined, of course, by the rate at which the metal dissolves, and this is proportional to the surface of the latter exposed in the cell. If 484 INORGANIC CHEMISTRY FOR COLLEGES the zinc poles of two cells of the same size and construction are connected and the copper poles are likewise connected, we can obtain from the combination twice the quantity of electricity in a given time that can be obtained from one of the cells; the com- bination delivers twice as many amperes as a single cell. By joining the two zinc poles we have doubled the area of the latter and the amount of the metal dissolved in a given time is doubled. Cells joined in this way are said to be connected in parallel. 570. The Intensity Factor of the Energy Transformed in Electrochemical Change. The fact has been mentioned a number of times that when two different substances are brought into con- tact a difference of potential is set up immediately as the result of the transfer of electrons from one to the other. This occurs when a metal is placed in a solution of one of its salts; in the case of zinc, for example, we have Zn 2 ^ Zn ++ . Some zinc atoms lose electrons and pass into solution as ions charged positively in respect to the metal, because the latter holds an excess of elec- trons and is, therefore, negatively charged. The ease with which this change takes place is determined in the case of any metal by the number of ions present in the solution; the more ions present the less the tendency of the metal to form ions, and the reverse is true. As a consequence, in studying the phenomenon quanti- tatively the concentration of the ions in the solution must be known. The metals differ from one another in the strength of the tendency they exhibit to form ions. Sodium, for example, passes readily into solution as ions when brought into contact with water; calcium forms ions less readily; in the case of zinc the tendency is much smaller. There is an analogy between the tendency of liquids to pass into vapor and the tendency of metals to pass into the forms of ions. In the case of the former the tendency is measured by the pressure of the vapor produced from the liquids; in the case of the latter the tendency is measured by what is called the electrolytic solution pressure. This pressure is the cause of the difference of potential set up when a metal is brought into contact with water or a solution of one of its salts. In order to measure the solution pressure of a metal it is only necessary, therefore, to determine this difference in potential. When a voltmeter is connected with the two poles of a cell, the ELECTROCHEMISTRY 483 instrument measures the difference between the potential of one pole and that of the other; the electromotive force observed is the difference between two potentials In the case of the Daniell cell, for example, the potential of the zinc pole is produced as the result of the passage of metallic zinc into ions, and that of the copper pole as the result of the change of copper ions to the metal. There is a difference in potential between the zinc pole and the solution of ions, and one between the copper pole and its ions. In order to distinguish between a difference of potential set up between a metal and a solution of its salt, and one set up between the two poles of a cell, the former type is called a single potential. The name is a satisfactory one because a single metal is involved in setting up the difference in potential, whereas the difference in potential of the two poles of a battery results from two potentials set up by two metals. When some metals are placed in normal solutions of their salts there is a greater tendency for the metal to pass into solution than for the ions to deposit as free metal; with certain metals the reverse is true. In the first case the metal becomes charged nega- tively and the solution positively, because some atoms pass into solution without electrons to form the ions and the excess of elec- trons left in the metal charge it negatively; the solution is posi- tively charged because it contains an excess of positive ions. In the second case the metal is charged positively and the solution negatively; this results from the fact that when the positively charged ions deposited on the metal they rendered the latter positive. The metals differ not only in regard to the positive or negative nature of the charge they assume when placed in contact with their solutions, but in the intensity of the charge; the difference in potential produced between a metal and a normal solution of its ions is definite and characteristic of the metal. The values for the single potentials produced in this way have been determined for all the metals. When the metal reacts with water at ordinary temperatures its single potential cannot be determined, but can be calculated from data obtained in other ways. The values of the single potentials set up when the metals are placed in normal solutions of their salts are given in the following table. This arrangement of the metals is known as the electromotive series. Fe(Fe++)., , + 0.14 Sb 8? Co. 05 Hg (Hg+) 1 02 Ni 05 Ag 1 08 Sn(Sn++)., , 0.13 Pd . 11? Pb (Pb ++ ). 0.15 Pt 1 2? H 2 28 Au (Au ') 1 4? Cu (Cu ++). - 0.60 486 INORGANIC CHEMISTRY FOR COLLEGES ELECTROMOTIVE SERIES OF THE METALS The figure after each metal is the potential in volts of a normal solution of a salt of the metal in contact with the metal. The figures in parentheses have not been determined directly; the values which are followed by an interrogation point are doubtful. K (+2.6) Cd + 0.15 Bi Na (+2.4) Ba (+2.4) Sr (+2.3) Ca (+1.9) Mg +1.5? Al +1.0? Mn +0.8 Zn +0.5 As -0.6? The signs + and in the table refer to the charge on the solu- tion. Down to cobalt the solutions are positively charged; this means that there is a greater tendency for the metal to form ions than for the ions to pass into the metallic condition. From cobalt to gold the greater tendency is in the reverse direction, and the solutions become negatively charged. In the case of hydrogen an electrode is formed by keeping saturated with the gas finely divided platinum supported on platinum foil. 671. In the above table the single potential of the hydrogen electrode is taken as 0.28 volt. The single potentials of the other elements have been determined by comparing them with that of hydrogen. No method has been devised of determining in an entirely satisfactory manner the single potential set up between an element and its ions. Consequently, there is doubt as to the correctness of the value assigned to hydrogen. On account of this fact, many chemists take the hydrogen electrode as the standard, and assign to its single potential the value 0. The values of the single potentials on this basis can be obtained by adding 0.28 to the numbers given in the table. The elements fall in the same order whatever standard is used; and it is this order which is of importance in interpreting the properties of the elements and their compounds. 572. The electromotive series of the metals has been used repeatedly throughout this book in interpreting the behavior of metals in a comparative way. We see now how it has been arrived at. The metals are arranged in the order of their ten- dency to pass into solutions as ions as measured by the intensity factor of the energy developed when such a change takes place. We have seen that the quantity factor of the electrical energy set ELECTROCHEMISTRY 487 free in this way is determined by the valence of the ion. When 1 gram-atomic-weight of magnesium, which has the valence 2, passes into solution as a magnesium ion, 2X96,500 coulombs of electricity are produced; the same quantity is set free when 1 gram-atomic-weight of iron passes into bivalent ions. In the first case 1.5 volts are developed, and in the second case but 0.14 volt. Since electrical energy is the product of its quantity and intensity factors, the chemical energy transformed into electrical energy in the case of magnesium is 2 X 96,500 X 1.5 = 289,500 joules and in the case of iron 2 X 96,500 X 0.14 = 27,020 joules. These facts make clear the difference in the energy content and activity of the two metals. The values of the potentials given in the table are those obtained when normal solutions of salts are in contact with the metals. At other concentrations the values are different. For example, in the reaction Zn - 2 ^Zn ++ , which is reversible, an equi- librium is set up as the result of the opposing tendency of the metal to pass into ions, and the tendency of the ions to deposit as metal. Increase in the number of metallic ions would, accordingly, bring about a new state of equilibrium, and the difference in potential set up between the metal and the solution would be less than before. In comparing the activity of two metals as measured by the potentials set up, false conclusions maybe arrived at if the concentrations of the ions of the metals are widely different. When a metal above hydrogen in the electromotive series reacts with an acid to form hydrogen, it dissolves because its tendency to form ions is greater than that of hydrogen ; the metal passes into solution and the hydrogen ions are discharged. As the metal dissolves and its ions increase in concentration, the potential becomes less and less but, in general, its value remains greater than that of hydrogen, and the gas continues to be evolved. 573. In the case of a metal below hydrogen, the gas is not ordinarily set free when it is placed in a solution of an acid, because hydrogen has a greater tendency to form ions than the metal under the conditions which exist. In this case there is an equilibrium set up, however, between the elements and their ions, as before, and if the metallic ions are removed as fast as they are formed the tendency of the metal to form ions must increase ; if it becomes greater than that of hydrogen the metal passes into solu- 488 INORGANIC CHEMISTRY FOR COLLEGES tion. This is what occurs when mercury is treated with concen- trated hydrobromic acid, and copper with concentrated hydro- chloric acid. In both cases compounds are formed between the acids and the halides produced, which yield an exceedingly small concentration of metallic ions. 574. Electromotive Series of Non-metallic Ions. The single potentials set up between certain non-metallic elements and their ions have been determined; some of these are given in the follow- ing table: I 2 (I~) -0.81 Cl 2 (Cr) -1.64 Br 2 (Br~) -1.36 O 2 (OH-) -1.94 2 (0~-)-1.49 When chlorine is in contact with water some of the element passes into solution as Cl~ ions; as a result, the solution becomes nega- tively charged in respect to the free element. The sign before the number expressing the difference in potential is, therefore, nega- tive, since in this table, as well as in that of the electromotive series of metals, the sign refers to the charge on the solution. It indicates in the case of all the elements listed above a tendency to pass into solution as negative ions. The numerical value of the difference in potential, disregarding the sign, is a measure of the tendency for ionization to take place the larger the number the greater this tendency. It is evident from an examination of the figures why free chlorine displaces iodine atoms from solutions: 2K + + 21- + C1 2 = 2K + + I 2 + 2Cr 675. The Electromotive Force of Electric Cells. When metals and solutions of salts are combined to make a cell, the electro- motive force of the combination is the result of two separate elec- tromotive forces produced in the cell; one is set up between the anode and the ions of the metal of which it is composed, and the other between the cathode and its ions. From the values given in the table on page 486 it is possible to calculate the electromotive force of a cell composed of two metals, each in contact with a normal solution of one of its salts. It will be recalled that the signs given in the table refer to the charge on the solutions and that those on the metals have the opposite sign. In the Daniell cell, for example, copper is charged positively and its potential ELECTROCHEMISTRY 489 with reference to copper sulphate is +0.60 volt; zinc is charged negatively and its potential is 0.5 volt. The difference in poten- tial between the two metals is evidently the difference between these two values, 0.60 (0.5) = 1.1 volts. If a cell is made up of iron and zinc the difference in potential is 0.14 ( 0.5) = 0.36 volt. In order to get as high an electromotive force as possible elements widely separated in the electromotive series are selected. It is found in practice that zinc on account of its cheapness is the best metal to use high in the series. Either copper or carbon with manganese dioxide are the most available substances from which to construct cathodes. When manganese dioxide is used, the reaction at the cathode which produces an electromotive force is the reduction of MnO2 to M^Oa, the oxygen reacting with hydrogen to form OH~ ions. The E.M.F. of a dry cell is approx- imately 1.6 volts. In but a few cases it is possible to construct a cell in which all the chemical energy set free is converted into electrical energy; in most cases there is a heat change, and when this occurs the elec- tromotive force is not that calculated in the way indicated above. This difference in the case of cells commonly used is small, however, and the change in potential due to differences in concentration of the ions is likewise relatively small for com- paratively great changes in concentration. The calculated values approximate, therefore, the observed ones. 576. Electrolytic Conduction. It has been pointed out that when an electric current flows through a wire there is no chemical change in the material of which the wire is made; as the current passes an atom gives up an electron to the next atom and receives another in turn, which it passes on; a constant stream is thus produced if a supply of electrons is furnished at one end of the wire and they are removed continuously at the other end. The conduction in a solution where the current is carried by the ions is pictured differently. In this case where the conduction is said to be electrolytic, the electrolyte, which is the acid, base, or salt in solution, undergoes decomposition at the electrodes. The nega- tive ions give up electrons at one pole and the positive ions take up electrons at the other. In the case of hydrochloric acid, for example, each chlorine ion loses an electron and is thereby con- 490 INORGANIC CHEMISTRY FOR COLLEGES verted into atomic chlorine, which changes immediately into chlorine gas; likewise the hydrogen ions take up an electrons x and become hydrogen atoms and then hydrogen gas. When the current has passed a short time the ions in the immediate vicinity of the electrode have been dis- the , Sit FIG. 35. Fre& charged, but tne process con- tinues because the charged elec- trodes attract to themselves the charged ions in the solution. There is thus set up a continuous stream of negative ions toward the pole charged positively and one of positive ions toward the negative pole; the positive ions travel in the direction of the flow of the positive current and the negative ions in the direction of the flow of the electrons or negative current. These conditions are shown in the accompanying diagram (Fig. 35). 577. Storage Batteries. In the type of cell which has been described, electrical energy is produced from chemical energy as the result of the change of a metal into ions and the reverse. The current in the Daniell cell is set up as the result of changes which may be indicated thus : Zn = Zn ++ + 2 and Cu ++ + 2 = Cu Zinc passes into solution at one pole and copper is deposited at the other, both changes liberating energy which sets up a flow of electrons from the zinc pole, where they were set free, through the connecting wire to the copper pole, where they were taken up, and, as a consequence, an electric current is produced. Both reac- tions are reversible, and if electrons are furnished from outside the cell and made to flow in a direction opposite to that in which they travel when the cell is producing electricity, the chemical changes effected will be reversed. When the cell is producing electricity the reaction at the zinc pole is Zn = Zn ++ + 2 ; when a supply of electrons is furnished by an outside current the reaction is Zn ++ + 2 = Zn, and a similar reversal takes place/ at ELECTROCHEMISTRY 491 the copper pole. It is evident, therefore, that we can obtain electricity from the Daniel! cell, or, by passing electricity into it, bring about chemical changes which lead to the separation of sub- stances from which electricity can be produced; in the latter case electrical energy is transformed into chemical energy, which later can be reconverted into electricity. When the Daniell cell is used in the first way it is said to be a primary cell; when used in the second way it is said to be a secondary or storage cell. For many reasons the Daniell cell is not economical when used as a storage cell, and other combinations are employed. In all cases, substances are selected which react reversibly without appre- ciable loss of energy as heat, and without the production of gases which escape from the cell. There are at present two important kinds of storage cells. 578. In the lead accumulator, when it is charged, one pole is lead dioxide, PbO 2 , the other metallic lead, and the electrolyte is sulphuric acid. When the poles are connected so that a current can flow, lead sulphate is formed at both the anode and cathode. Metallic lead reacts with the sulphate ion to form lead sulphate, Pb + SO 4 ~ ~ = PbSO 4 + 20 and the electrons on the negative ion are set free. The hydrogen ions at the other pole take up electrons and are thus converted into atoms which reduce the dioxide to oxide; and the latter is converted by the acid present into lead sulphate and water : 2H + + 20= [2H] [2H] + PbO 2 = H 2 O + [PbO] [PbO] + H 2 S0 4 = PbSO 4 + H 2 2H + + PbO 2 + H 2 SO 4 + 20 = 2H 2 O + PbSO 4 The current is, thus, produced as the result of the liberation of 2 electrons from the SO 4 radical at one pole and the transfer of 2 electrons to 2H + ions at the other. The complete equation for the reaction which takes place during the discharge of the cell is, accordingly, PbO 2 + Pb + 2H 2 SO 4 = 2PbS0 4 + 2H 2 + Energy 492 INORGANIC CHEMISTRY FOR COLLEGES When a current is passed into the cell to charge it, the reaction at one pole is PbSO 4 At the other pole an SO 4 ion is taken up to form a persulphate which is immediately hydrolyzed by water: PbSO 4 + SO 4 ~~ = [Pb(SO 4 ) 2 J + 20 [Pb(SO 4 ) 2 ] + 2H 2 O = PbO 2 + 2H 2 SO 4 PbSO 4 + SO 4 + 2H 2 O = PbO 2 + 2H 2 SO 4 + 20 The summation of the reactions at the two poles leads to an equation identical with that which represents the discharge of the cell, except that it is written in the reverse direction : charge Energy + 2PbS0 4 + 2H 2 O ^ PbO 2 + Pb + 2H 2 SO 4 discharge It is not difficult to remember in which direction the equation is to be read to express either discharge or charge, when it is recalled that the formation of a salt and water from an acid and a metal or an oxide is associated with the setting free of energy. Under ordinary circumstances this energy appears as heat, but by the arrangement used in a storage cell it is given off in the form of an electric current. The cause of the setting up of the current can be expressed very simply; it is produced as the result of the change of one atom of metallic lead, which has valence, Pb, and of 1 atom of lead as the dioxide with the valence 4, Pb IV , into 2 atoms of lead in the form of the sulphate in which the metal has the valence 2,Pb + Pb IV = 2Pb n . The metallic lead loses 2 electrons, Pb = Pb + +20, which convert the quadrivalent lead to bivalent lead, 2 + Pb + + + + = Pb + + ; it is the transfer of these electrons in the circuit which causes the current. The poles of the lead accumulator are made in the form of flat grids of lead the interstices of which contain the active materials of the cell, finely divided lead in one case and lead dioxide in the other. The electromotive force produced by each cell is about 2 volts. When a cell is in the charged condition the solution contains sul- phuric acid; as it is discharged lead sulphate is deposited on both poles and water is formed. It is possible, therefore, by an observa- ELECTROCHEMISTRY 493 tion of the specific gravity of the solution, which is a measure of the amount of acid present, to tell to what extent the cell has been discharged. 579. The lead accumulator is very heavy, since it is made of lead and involves the use of sulphuric acid, and, consequently, many attempts have been made to utilize other reactions in the construction of storage cells. The most successful of these is known as the Edison cell, in which the poles, when charged, are nickelic oxide, Ni 2 O3, and iron. The electrolyte is a solution of potassium hydroxide. When the cell discharges the nickelic oxide is changed to nickelous hydroxide, Ni(OH)2> and the metallic iron to ferrous hydroxide : discharge Fe + Ni 2 O 3 + 3H 2 O ^ Fe(OH) 2 + 2Ni(OH) 2 + Energy charge The changes in valence which cause the current are represented as follows : discharge Fe + 2Ni m ^ Fe 11 + 2Ni n + Energy charge Iron is above nickel in the electromotive series and, as a conse- quence, the energy set free when it is oxidized, Fe * Fe 11 , is greater than that required to reduce an equivalent amount of nickel, 2Ni ni * 2Ni n ; the conversion in this direction thus liberates energy which appears in the form of electricity. 580. Uses of the Electric Current in Chemistry, (a) For Heating. All substances offer more or less resistance to the flow of an electric current through them, and, as a consequence, a part of the electrical energy is changed into heat. When it is desired to avoid loss of energy in this way, as in the conduction of elec- trical energy from one place to another, a conductor is chosen which has a low resistance; copper is ordinarily used for this purpose (see table, page 443). When, however, it is desired to convert electrical energy into heat, a conductor which has a high resistance is selected, and for this reason alloys, carbon, and non- metallic substances, are commonly used. The heat generated in a circuit is equal to the square of the current multiplied by the resistance. When the current is 1 ampere and the resistance is 1 ohm the heat developed is 0.2388 calorie ( = 1 joule per second). 494 INORGANIC CHEMISTRY FOR COLLEGES When comparatively small quantities of heat are to be devel- oped, a current is sent through a wire in order to deliver the heat where it is required. For small furnaces tubes of fused silica or porcelain are wound with a wire made of an alloy which has a high resistance and a high melting-point, and does not oxidize rapidly in air at high temperatures. An alloy containing 65 per cent chromium and 35 per cent nickel is well adapted to this pur- pose^commercial " nichrome " or " chromel " wire contains these elements and usually some iron. In order to prevent loss of heat the furnace is covered with asbestos, which is a poor con- ductor. For technical use where very high temperatures are required, the resistor is usually powdered coke, which offers a great resistance to the current. A number of the products manufactured in the electric furnace are made from carbon and other substances, and the mixture used is itself the resistor. Calcium carbide is made in this way from lime and carbon (217). For high-temperature work on a small scale a so-called arc- furnace may be used. In this case the material is placed in a cru- cible of graphite, which serves as one electrode; when a carbon pole is brought into contact with the crucible and then separated from it, an arc is produced and the crucible is heated to a very high tem- perature as long as the arc is maintained. So-called induction furnaces are used in the arts to melt iron and other metals. 581. (b) Use of the Electric Current in Preparing Elements and Compounds. The more active of the metals are prepared by the electrolysis of their compounds, on account of the fact that their oxides are not reducible by carbon at the temperature produced when the latter unites with oxygen. Sodium, potassium, calcium, magnesium, and aluminium are obtained in this way. The details of the methods used will be described later. The active non-metallic elements, chlorine and fluorine, are also obtained by electrolysis, as we have seen. Sodium hydroxide and the hydroxides of the other alkali metals are prepared industrially from salts by electrolytic methods. 582. (c) Use of the Electric Current in the Purification of Metals. The presence of small amounts of foreign substances in metals ELECTROCHEMISTRY 495 often affects materially their properties. The electrical conduc- tivity of copper, for example, is markedly affected in this way, and, as a consequence, it is brought to a high state of purity before being used for electrical 'purposes. This is most readily accomplished by depositing the metal from a solution of its salt by means of an electric current. The crude copper as it comes from the smelter is suspended in a bath containing a solution of copper sulphate, and serves as one pole; the other is a thin plate of the pure metal. The poles are connected with a source of electricity in such a way that the metal will be deposited on the pure copper. The copper ions are deposited on the negative pole and copper passes into solution at the positive pole. If the crude copper contains any metal above copper in the electromotive series, such as zinc, it will pass into solution; elements less active than copper, such as gold and silver, will not be converted into ions but will fall to the bottom of the tank along with compounds of copper with non-metals, such as the sulphide and arsenide. At the cathode the metals more active than copper will not be deposited, because the tendenc} r of copper ions to. lose their charge and pass into the metallic condi- tion is greater than that of the elements above it in the electro- motive series. As far as the copper is concerned, the reaction at one pole is just the reverse of that at the other; the energy required to set copper free from its ions at the cathode is equal to that lib- erated by the ionization of the metal at the anode. Consequently, the energy required is only that necessary to overcome the resistance of the flow of the current through the solution, which results in the movement of the ions. This is reduced by keeping the liquid in the cell well stirred and having the plates placed close together. In practice the electrolysis is carried out in a trough lined with lead into which are placed alternately the anodes and cathodes. Under these conditions an electromotive force of less than 0.5 volt is required. The copper obtained in this way is about 99.8 per cent pure. Very large amounts of gold and silver are recovered from the sludge obtained from the tank. The process is also the chief source of tellurium, which is obtained from the copper telluride present in the crude copper. Iron which is to be used in the cores of electro-magnets must be as free as possible from impurities. The metal can be purified 496 INORGANIC CHEMISTRY FOR COLLEGES for this purpose by a method analogous to that used in the case of copper, and when the highest purity is desired this method is employed. 583. Electroplating. This process can be carried out in a manner analogous to that described for the purification of copper. The cathode is the object to be coated, the anode the metal to be used, and the electrolyte a solution which furnishes the ions of the metal at such a concentration that when they deposit, a closely adhering layer of the metal is formed. Silver and gold are commonly used in electroplating. As these elements are very inactive, they are deposited from solutions of most of their salts when brought into contact with more active metals. If iron, for example, is placed in a solution of silver nitrate, the metal goes into solution and silver is deposited in the form of a loose powder on the surface of the metal. It is evident, therefore, that silver nitrate could not be used as the electrolyte in electroplating silver, for the metal would be deposited chemically in a non-adhering form. A salt must be used, accordingly, which produces such a low concentration of silver ions, that the ten- dency of the latter to pass out of solution in the presence of the more active metal is reduced to the point where silver does not separate; the double cyanide of silver and potassium, KAg(CN)2, is such a salt. When this compound is dissolved in water the ions formed are chiefly K + and Ag(CN)2~, but a very small concen- tration of silver ions is produced as the result of the further ionization of the negative Ag(CN)2~ ion into Ag + and 2CN~. The concentration of the silver ion is so small that its direct deposition on the metal without the use of the current does not take place. Even under these conditions it is difficult to deposit silver electrolytically on iron and the metal is usually first coated with copper upon which the silver is deposited. Double cyanides are also used in electroplating with gold and with nickel. In making electrotypes for use in a printing press, a wax mold is first made with the impression desired. Its surface is next coated with finely powdered graphite to make it a conductor of electricity. It is then suspended in a bath of a copper salt and connected with the negative pole of a supply of electricity. A plate of copper in the bath is connected with the positive pole. When the deposit of copper on the mold is sufficiently thick, the metal is stripped ELECTRO-CHEMISTRY 497 from the wax, and backed with stereotype metal, and finally mounted on a wooden block. 584. Oxidation and reduction are carried out by means of the electric current, because it is possible to obtain hydrogen and oxygen when aqueous solutions are electrolyzed. It has been stated, for example, that perchloric acid, HCICU, is now made by the electrolytic oxidation of hydrochloric acid, and potassium chlorate in an analogous way from potassium chloride. Elec- trolytic oxidation and reduction are conveniently applied in the manufacture of certain important organic compounds. EXERCISES 1. The reactions Zn^Zn ++ -f20 and H 2 ^ 2H+ -f 2 are reversible. Show how the facts in regard to these changes are in accord with the law of mobile equilibrium. 2. Draw a diagram indicating the flow of the electric current and the changes which take place when zinc is in contact with copper in the presence of a solution of hydrochloric acid. 3. Draw a diagram to illustrate how an experiment could be carried out to test Faraday's law. 4. What weights of copper and chlorine will be liberated when (a) 20,000 coulombs are passed through a solution of copper chloride, and (6) when a current of 5 arnperes is passed for three hours through the solution. 5. An electric current was passed first through a solution of a silver salt and then through one of a copper salt in which the metal had the valence 2. It was found that in the same time 2.3733 grams Ag and 0.6992 gram Cu were deposited. Assuming 107.88 as the atomic weight of silver calculate the atomic weight of copper. 6. Complete the following expressions, replacing a, 6, c, etc., by the proper units: (a) number of coulombs = aX&; (6) number of joules cXd; (c) number of kilowatt-hours = eXf=gXhXi. 7. It requires 2.3 volts to liberate chlorine from a solution of sodium chloride, (a) Calculate the number of joules required to set free 1 gram- atomic-weight of chlorine: (6) how many kilowatt-hours is this number of joules? (c) If electrical energy costs 10 cents per kilowatt-hour, what is the cost of the electrical energy required to set free 1 kilo of chlorine? 8. What does the electrical energy cost to run a 60-candle-power tungsten incandescent lamp for 1 hour if the energy cost 10 cents per kilowatt-hour and the lamp consumes 1 . 2 watts per candle power? 9. Calculate the E.M.F. of cells made up of the following pah's of elements in contact with normal solutions of the respective salts, and state in which direction the positive current would flow in each case through a wire connect- ing the two metallic poles: (a) Al and Zn, (6) Zn and Co, (c) Zn and Pb, (d) ZnandAu, (e) Hg and Ag. 498 INORGANIC CHEMISTRY FOR COLLEGES 10. Explain how you can tell by the determination of the specific gravity of the sulphuric acid in a lead storage cell whether or not it is charged. 11. The resistance per foot of number 20 nichrome wire, which is 0.81 mm. in diameter, is 0.52 ohm. (a) How many feet of the wire are required to furnish a resistance of 100 ohms? (6) If the difference of potential at the ends of a wire having a resistance of 100 ohms is 110 volts how many cou- lombs pass through the wire per second (current = E.M.F. -r- resistance)? (c) How many calories will be obtained if the current passes for one hour? 12. Calculate the cost of 1,000,000 calories produced by (a) burning coal at $10 per ton (1 gram coal yields 8000 calories) and by (6) the electric current at 2 cents per kilowatt-hour. 13. Two dry cells were tested by applying the terminals of a voltmeter to the zinc and carbon poles of the cell. In each case the meter registered 1.9 volts. When tested with an ammeter each cell registered 30 amperes, (a) If the carbon poles of the two cells are connected to each other and the zinc poles are connected, what would the voltmeter and the ammeter register if joined to the carbon and the zinc poles? (6) If the carbon pole of one cell is connected with the zinc pole of the other and the measuring instruments connected with the free poles what would the readings be in each case? (c) If six partially used cells were available which, when tested, showed 1.6 volts and 8 amperes how should they be connected to use with a small incan- descent lamp requiring 3 volts? 14. When the salt of the formula KAg(CN)2 is used in electroplating why is potassium not deposited at the cathode? 15. How could gold be separated from zinc by electrolysis? CHAPTER XXXIV THE PROPERTIES OF OXIDES, HYDROXIDES, AND SALTS 585. A very large number of compounds are known which contain metallic atoms. With the aid of the electromotive series and the periodic law, however, it is possible to systematize the facts to be presented, by grouping the compounds into several classes, the members of which exhibit similar properties that vary progressively with the atomic weight and activity of the metal. In this chapter the properties of compounds and their relationships will be considered in a general way, and later specific facts will be given in regard to those particular compounds which are of interest from a theoretical point of view or have found important practical applications. OXIDES 586. All the metals form well characterized oxides, and in the case of most metals two or more compounds of this class are kriown. The highest valence which can be shown by a metal toward oxygen is indicated by its position in the periodic classification, and in most cases oxides are known in which this valence is shown; the most striking exceptions occur in the case of the metals in the eighth group, where osmium and ruthenium alone form oxides in which the metals have the valence 8. Oxides are formed by heating metals in the air or oxygen, or by decomposing salts of the metals. Whether or not a metal burns in the air is determined in certain cases by its physical con- dition. Under ordinary circumstances the metals in the elec- tromotive series down to and including zinc burn in the air; iron and lead can be obtained in such a finely divided condition that they take fire spontaneously when brought into the air. The metals down to silver will burn in oxygen. The composition of 499 500 INORGANIC CHEMISTRY FOR COLLEGES the oxide formed when a metal burns is determined by the activity of the latter and its valence. Sodium and potassium form per- oxides, Na202 and K^O*, which contain an excess of oxygen over that present in the normal oxides, Na 2 O and K 2 O. The alkaline earths, magnesium, zinc, cobalt, nickel, tin, lead, and copper give oxides of the general formula MO; aluminium, chromium, arsenic, antimony, and bismuth, oxides of the formula M^Os; and man- ganese and iron, oxides of the formula MsCX in which two atoms of the metal have the valence 3 and one the valence 2. 587. The chemical properties of an oxide of a metal are determined by the activity of the metal and its valence. Some oxides react with acids to form salts, some with bases to form salts, and some in both ways; other oxides react with neither acids nor bases. The oxides of the more active metals the alkali metals and those of the alkaline earths dissolve in water, and form strong bases, which react with acids to form salts. Mag- nesium oxide is also basic. Aluminium oxide dissolves in both acids and bases, and, therefore, exhibits both acidic and basic properties. Certain metals which can show a high valence form oxides in which they function as acid-forming elements. Man- ganese, for example, which is in the seventh group in the periodic classification, forms an oxide of the formula M^Oy, which is the anhydride of permanganic acid, HMnO-i; the latter is a strong acid and forms salts which resemble in composition and properties those derived from perchloric acid, HCICU. As has already been emphasized, increase in valence of an element toward oxygen is associated with the development of acidic properties; and this occurs even when the element is a metal. Manganese is like iron in many of its physical and chemical properties. When it has the valence 2 it acts as a strong base-forming element, with the valence 3 it is still base-forming but less active. The temperature at which the oxides are reduced by hydrogen or carbon varies with the activity of the metal and the valence it shows in the oxide. When heated with hydrogen the oxides down to and including manganese can be reduced to lower oxides, if they exist, but complete reduction to the metal cannot be effected. The oxides of cadmium and the metals below it can be easily reduced to the metal by hydrogen. The oxides of mer- cury and the less active metals are readily decomposed by heat THE PROPERTIES OF OXIDES, HYDROXIDES, AND SALTS 501 alone. At the temperature obtainable in a blast-furnace, the oxides of the metals from manganese down can be reduced by carbon. HYDROXIDES 588. The hydroxides of the alkali metals are soluble in water and are formed by the action of water on the oxides of the metals; they are prepared by the electrolysis of the chlorides. The hydroxides of the alkaline earths are difficultly soluble in water and are made by the action of water on the oxides. The hydrox- ides of the other metals, being insoluble, are formed as precipi- tates when the salts of the metals are treated with a solution of a soluble hydroxide. When, however, a soluble base is added to a solution of a salt of silver, mercury, cuprous copper Cu 1 , or aurous gold Au 1 , an insoluble oxide is formed. In the case of silver, some of the oxide dissolves as hydroxide. All the hydroxides except those of the alkali metals are decom- posed into oxides and water when heated. They all dissolve in acids to form salts. The hydroxides of univalent metals are strong bases, and, as a result, form neutral salts with strong acids that are not hydrolyzed by water. Cuprous salts and aurous salts are hydrolyzed. The hydroxides of the bivalent metals are relatively strong bases and their salts are hydrolyzed but slightly if at all at ordinary temperatures. The hydroxides of the tri- valent metals are weak bases; their salts with strong acids are appreciably hydrolyzed and those with weak acids, such as car- bonic acid and hydrogen sulphide, are completely hydrolyzed by water. 589. Certain hydroxides which are precipitated by a solution of sodium hydroxide are not precipitated by one of ammonium hydroxide if ammonium salts are present. This fact is utilized in analytical chemistry in quantitative separations of metals and is, therefore, of importance. The explanation is based on a con- sideration of the facts from the point of view of the law of molecu- lar concentration. When a hydroxide is precipitated from solution, the com- pound is formed as the result of the union of the metallic and the hydroxyl ions. For example, in the case of a bivalent metal the reaction may be indicated as follows: M + + +20H~^M(OH)2. 502 INORGANIC CHEMISTRY FOR COLLEGES If the solubility of the hydroxide is very small, the amount in solution is completely ionized and an equilibrium is set up between the solid present and its ions. According to the law of molecular concentration, in the case of the hydroxide of a bivalent metal the concentration of the metallic ions multiplied by the square of the concentration of the hydroxyl ions is a constant. Similar equilibria exist between the ions of other slightly soluble com- pounds; the constant for such an equilibrium is commonly called the solubility product. If the ions of a slightly soluble compound are brought together in solution, they will react to form the com- pound, which will precipitate until the solubility product of the ions is reached. For example, when sodium hydroxide, which furnishes a high concentration of OH~ ions ; is added to a strong solution of calcium chloride, calcium hydroxide is precipitated because the solution furnishes a higher concentration of OH" and Ca ++ ions than that indicated by the solubility product of these ions in equilibrium with calcium hydroxide. When ammonium hydroxide, which is a weak base, is added to a solution of a cal- cium salt, the solubility product is not exceeded and the calcium hydroxide is not precipitated. Ammonium hydroxide does pre- cipitate magnesium hydroxide, however; the solubility of the latter is much less than that of calcium hydroxide, and, as a result, the solubility product of magnesium and hydroxyl ions is less. When ammonium hydroxide is added to a solution of a mag- nesium salt, magnesium hydroxide is precipitated until the concen- trations of the OH~ and Mg ++ ions are reduced to those indicated by the solubility product. If a solution containing ammonium hydroxide and ammonium chloride is added to a magnesium salt, the hydroxide is not precipitated. This is due to the fact that the extent of ionization of ammonium hydroxide, and, therefore, the concentration of hydroxyl ions in its solution, is reduced in the presence of ammonium salts. There is an equilibrium established when ammonium hydroxide ionizes: If ammonium ions, produced from ammonium chloride, NH 4 C1 = THE PROPERTIES OF OXIDES, HYDROXIDES, AND SALTS 503 are added to the solution of ammonium hydroxide, the equilibrium in the latter case is disturbed and shifts from right to left; as a result, NH 4 + and OH~ unite to form NH 4 OH, and the concen- tration of the OH~ ions is reduced. If a sufficient quantity of an ammonium salt is added to a solution of ammonium hydrox- ide, the concentration of the OH~ ions is reduced to such an extent that in the presence of a magnesium salt the solubility product of magnesium and hydroxyl ions is not obtained and the precipi- tation of magnesium hydroxide does not take place. The precipitation of ferric hydroxide, Fe(OH)a, by ammonium hydroxide is not prevented by the presence of ammonium salts because the solubility of the metallic hydroxide is very small, and it furnishes, therefore, a very low concentration of hydroxyl ions. It is impossible by the addition of ammonium salts to reduce the concentration of hydroxyl ions in a solution of ammonium hydroxide to such an extent that the solubility prod- uct of ferric and hydroxyl ions is not exceeded; precipitation of ferric hydroxide, accordingly, occurs. Ferrous hydroxide, on the other hand, is more soluble, and its precipitation by ammonium hydroxide is prevented by ammonium salts. The solubilities of the hydroxides of most of the metals having the valence 2 are such that they are not precipitated by ammonium hydroxide in the presence of ammonium salts; the hydroxides of the trivalent metals are much less soluble and, as a consequence, they are precipitated under these conditions. SALTS 590. The methods of preparing salts (142, 148) and some of the properties which they possess in general (241, 242, 246, 249, 251) have already been described at some length. The solubilities of the salts of the more important acids which have been con- sidered, have been stated and the action of heat on acid, basic, and neutral salts has been emphasized (204, 282, 294, 307, 367, 373, 408, 428, 460, 462, 511). 591. Isomorphism. When a salt crystallizes, the form it assumes is characteristic of it. While the size and general shape of the crystal may vary according to the conditions under which crystallization takes place, the angles at which the faces cut one 504 INORGANIC CHEMISTRY FOR COLLEGES another are fixed. In general, on the slow evaporation of a solu- tion containing two or more salts the latter will be deposited as individual crystals. In the case of certain salts, however, which are closely alike in crystalline form, crystals of but one kind will be produced, which contain the salts in the proportion in which they were present in the solution. Salts which produce such mixed crystals are said to be isomorphous. In general, isomorphous salts are closely related in composition; for example, zinc sul- phate, ZnSC>4,7H2O, and manganese sulphate, MnSO4,7H2O, are isomorphous, and it is, therefore, impossible to separate them by crystallization. 592. Effect of Temperature on the Solubility of Salts. The change in the solubility of salts with change in temperature is a fact of great importance, which is utilized in separating salts and in their purification. Crytallization is one of the most effective processes used in the preparation of soluble chemical compounds. The change in solubility with rise in temperature is markedly different in the case of different salts; for example, under these conditions the solubility of potassium nitrate increases greatly, that of sodium chloride only slightly, whereas anhydrous sodium sulphate is more soluble in cold than in hot water. These differ- ences in the effect of increased temperature on the solubility of salts can be traced to the energy changes which take place when solution occurs. When potassium nitrate is dissolved in water the temperature of the solution falls far below that of the solvent; much heat is absorbed from the surroundings: Solid KNO 3 + water = solution of KNO 3 - 8500 cal. This equation, which expresses the heat of solution of potassium nitrate, may be abbreviated to read, KNOs+Aq = 8500 cal. It means that when 1 gram-molecular-weight of the salt is dissolved in such an amount of water (aqua) that further addition of the solvent produces no heat change, the heat absorbed is 8500 calories. There is an equilibrium between the solid and the solution, and since the transfer of the salt in either direction is accompanied by a heat change, the law of van't Hoff must apply. It will be recalled that rise in temperature shifts the equilibrium so that THE PROPERTIES OF OXIDES, HYDROXIDES, AND SALTS 505 heat is absorbed; as a consequence, under these conditions the salt becomes more soluble. In the case of sodium chloride the change is in the same direc- tion as with potassium nitrate, for NaCl+Aq = -1200 cal; but the heat change is much smaller and the effect of rising tempera- ture is, as a result, less. Anhydrous sodium sulphate dissolves in water with rise in temperature: NaoSCU+Aq = +460 cal. In accordance with the law of mobile equilibrium the solubility of the salt decreases with rise in temperature. Other changes than solution occur when certain salts dissolve in water. Among these are hydration, hydrolysis, and ionization; and as these processes produce heat effects, the change in solubility of a salt with change in temperature cannot always be determined by a consideration of its heat of solution. In the diagram on page 506 are plotted the solubility curves of a few salts. The concentrations are expressed as the number of grams of the anhydrous salt which are held in solution by 100 grams of water. The slope of the curve in each case indicates the change in solubility with rise in temperature; that of sodium chloride is nearly horizontal, whereas the curve of potassium rises steeply. Advantage is taken of the fact that salts absorb heat when they pass into solution, in making freezing mixtures for low tempera- ture work. A few examples of such mixtures in which the ingre- dients are expressed as parts by weight are as follows: 30 NEUCl and 100 water at 13.3, temperature reduced to -5.1; 60 NH 4 N0 3 and 100 water at 13.6, temperature reduced to -13.6; 110 Na 2 S 2 O3,5H 2 O and 100 water at 10.7, temperature reduced to -8; 250 CaCl 2 ,6H 2 O and 100 water at 10.8, temperature reduced to -12.4. When the effect of increased temperature on the solubility of two salts is markedly different, it is easy to separate them from each other by crystallization. A case of this kind is potassium nitrate and sodium chloride. If water is saturated with the two salts, say at 60, and then cooled, to say 10, but a small quantity ' of sodium chloride will crystallize out, since its solubility at 60 is about 38 grams in 100 of water and about 36 grams at 10. In the case of potassium nitrate the solubilities at these tem- peratures are 110 and 17 grams, respectively, and a large part of the salt will separate when the solution cools. If the mixture 506 INORGANIC CHEMISTRY FOR COLLEGES of the two salts obtained is recrystallized from hot water, the potassium nitrate which separates is practically free from salt. It is seen from the diagram that at room-temperatures potas- sium nitrate is much less soluble than sodium nitrate. Conse- 30 40 50 60 70 Temperature ,Deg. Cent. FIG. 36. quently, if boiling water is saturated with sodium nitrate and potassium chloride and the solution allowed to cool, a double decomposition takes place and potassium nitrate crystallizes out. This fact is utilized in manufacturing potassium nitrate from Chile saltpeter. In the great salt beds at Stassfurt and other places the THE PROPERTIES OF OXIDES, HYDROXIDES, AND SALTS 507 salts are found in the mines deposited in layers, the positions of which are in accord with the solubilities of the several salts. This fact shows that they were deposited as the result of the evaporation of an aqueous solution. 593. The Solubility of Salts in Acids. The difficultly soluble salts of weak acids are more or less soluble in strong acids, for example, calcium phosphate, which is practically insoluble in water, dissolves in nitric acid. The explanation of this fact is based on a consideration of the reactions which take place between the ions present in the solution. When calcium phosphate is in contact with water a very small amount of the salt passes into solution and forms ions: + 2PO 4 --- Nitric acid furnishes ions also : HN0 3 ^ H + + N0 3 " When phosphoric acid is in solution it ionizes almost wholly as follows : ^H + -f H 2 P0 4 ~ Just as hydrogen ions and hydroxyl ions when brought together unite to form water H + + OH~ + HoO so hydrogen ions unite with P0 4 ~" ~ ions to form the ion This occurs when calcium phosphate is placed in nitric acid; the removal of the PO 4 ~ " ions disturbs the equilibrium between the solid calcium phosphate and its solution, and more of the salt passes into solution until it finally dissolves. The ions finally present are Ca ++ ; NO 3 ~, H 2 PO 4 ~, and H + . In general, a salt of a weak acid will dissolve in a strong acid as the result of the fact that the hydrogen ions furnished by the strong acid unite with the radical of the weak acid to form the undissociated free acid or, as in the case of phosphoric acid, an undissociated ion. The cause of the change is the same as that which leads to the neutralization of a base by an acid. If in any way the ions of a very slightly dissociated substance are brought together, they will unite until their concentrations are those which 508 INORGANIC CHEMISTRY FOR COLLEGES can exist in equilibrium with the undissociated substance, that is, until the solubility product is reached. Whether or not the salt of a weak acid will dissolve in a strong acid is determined not only by the relative strength of the acids concerned, but also by the solubility of the salt in water. This can be shown by a consideration of the sulphides of zinc and copper. In the case of the former a certain amount of the salt passes into solution and ionizes: ZnS^Zn + + + S- When hydrochloric acid is added, hydrogen ions are present and an equilibrium is set up between them and the sulphur ions: 2H + + S" If the concentration of the S~~ ions furnished by the zinc sul- phide is greater than that of these ions when in equilibrium with H^~ ions, the equilibrium will shift and H^S molecules will be formed. More zinc sulphide will pass into solution as the result of the removal of S ions, and finally the salt will dissolve. In the case of copper, the sulphide is so little soluble that the S~~ ions are not removed, and the sulphide does not dissolve. It is for the reasons just stated that hydrogen sulphide will precipitate copper sulphide from a solution of one of its salts in the presence of hydrochloric acid, but will not precipitate zinc sulphide under the same conditions. Zinc sulphide is precipitated, however, by ammonium sulphide, because the solution contains no hydrogen ions. 594. The Use of Sulphides in Qualitative Analysis. The difference in the solubility of the sulphides in water, and in acids, is utilized in qualitative analysis. The sulphides of the alkali metals are soluble in water, and those of the alkaline earths and magnesium are hydrolyzed by water; they are not, as a conse- quence, precipitated by hydrogen sulphide. The sulphides of the metals in the electromotive series from magnesium to iron, inclu- sive, when they are in the bivalent condition, are soluble in dilute acids; they are precipitated by ammonium sulphide but not by hydrogen sulphide if an acid is present. The sulphides of the THE PROPERTIES OF OXIDES, HYDROXIDES, AND SALTS 509 metals from cobalt to gold are insoluble in dilute acids and are precipitated by hydrogen sulphide in the presence of acids. It is possible, therefore, to separate the metals into three large groups which can be divided into subgroups (Appendix VI), and the process can be continued until a complete separation has been effected. The method of doing this is studied in detail in quali- tative analysis. 595. The Solubility of Salts in a Solution of Ammonia. Com- pounds of certain metals which are insoluble in water dissolve in a solution of ammonia. Among the more important of these are the compounds of silver, zinc, copper, nickel, and cobalt. In all cases ammonia unites with the metallic atom and substances more or less soluble in water are formed. Silver chloride, for example, is converted into a compound of the composition AgfNHs^Cl, which gives the ions Ag(NH3)2 + and Cl~, and is readily soluble in water. In the case of zinc 4 molecules of ammonia are added to the metallic ion. Zinc hydroxide, which is insoluble in water, dis- solves in ammonia as the result of the formation of the compound Zn(NH3)4(OH)2. Cuprous copper gives the ion Cu(NH3)2 + , and cupric copper the ion Cu(NH3)4 +4 ~. 596. The Effect on the Solubility of a Salt of a Substance Yielding an Ion of the Salt. When a salt is difficultly soluble in water there is an equilibrium set up between the solid and the saturated solution. For example, in the case of lead iodide the equilibrium can be represented thus, PbI 2 (solid) Pb ++ +2I-, since all of the salt in solution is ionized. It is evident, from an application of the law of molecular concentration, that if either lead ions or iodine ions are added to the solution the equilibrium must shift, and, as a result, the change which takes place is that indicated by reading the equation from right to left; lead iodide comes out of solution and the solubility of the salt is thereby reduced. The fact that the solubility of salts is decreased in this way is utilized in analytical chemistry because it is advisable to reduce as far as possible the solubility of difficultly soluble salts and thus effect quantitative separations. For example, in precipitating 510 INORGANIC CHEMISTRY FOR COLLEGES silver chloride in the quantitative determination of chlorine ions, a slight excess of the silver salt is used in the precipitation to reduce the solubility of the silver chloride formed. 597. The Solubility of Salts in Solutions of Other Salts. In certain cases salts which are insoluble in water are soluble in solu- tions of other salts. Silver cyanide, AgCN, for example, is insol- uble in water but dissolves in a solution of potassium cyanide, KCN, as the result of the formation of a compound of the composition KCN,AgCN. The formula is preferably written in the form KAg(CN)2 because, when it dissolves in water, the ions produced are K~*~ and Ag(CN)2~. There are many salts of this type which are formed as the result of the combination of two salts. 598. Double Salts. Many compounds consisting of two salts in combination are formed from soluble salts; if potassium sulphate and aluminium sulphate are dissolved in water and the mixture of the two solutions is allowed to evaporate, crystals of a double salt of the composition K2SO4,Al2 (864)3, 24H2O are formed; the compound is commonly called alum. Many double chlorides are known. The one of the composi- tion K^PtCle is formed when potassium chloride is added to a solu- tion of platinic chloride, PtCU. Since it is but slightly soluble in water it is used in the quantitative determination of potassium or of platinum. The compound is called potassium chloroplatinate, because it may be considered as derived from platinic acid, H^PtOs, by the replacement of 3 bivalent oxygen atoms by 6 univalent chlorine atoms. Chloroplatinic acid, H^PtClojOH^O, is formed from hydrochloric acid and platinic chloride. We have already seen that certain sulphides unite to form com- plex salts (478, 489) and that they bear a similar relation in com- position to the salts of oxygen acids; ammonium thioarsenate and ammonium arsenate have the composition represented by the formulas (NH4)aAsS4 and QTH^sAsQ*. Many complex salts of this type are known. Since in many of these one metal functions as an acid-forming element, it is clear why in the majority of cases complex salts are produced from a salt of a metal which can form acids a weakly metallic element and a salt of a metal which forms strong bases. The stability of the double salts when in contact with water varies markedly. Some, like potassium ferrocyanide, THE PROPERTIES OF OXIDES, HYDROXIDES, AND SALTS 511 Fe(CN)2,4KCN or K 4 Fe(CN)o, are stable and yield complex ions, while others, like alum, decompose in aqueous solution into their constituents and exist only in the solid condition. A few of the more important double salts will be described later. The compounds formed by the union of salts derived from different acids are called mixed salts; those which contain but one acid radical and decompose largely in solution into the simple salts of which they are composed are called double salts; and those derived from a single acid that are stable in solution and yield complex ions are called complex salts. 599. The lonization of Salts. Although the acids differ widely in the extent to which they ionize, the salts of most acids are highly dissociated in solution. For example, sulphuric acid is ionized to the extent of about 61 per cent in a one-tenth normal solution at 18, and carbonic acid even at a higher dilution, N/25, is but 0.2 per cent ionized; on the other hand, the ionization of potassium sulphate is 72 per cent and that of potassium carbonate about 70 per cent. The ionization of acetic acid and that of its sodium salt in N/20 solution are 2 per cent and 79 per cent, respectively. The extent to which a salt ionizes is determined largely by the valence of the ions it yields. Salts which yield two univalent ions are the most highly dissociated. The ionization values of a few salts of this type at 18 in one-tenth normal solution are as follows: sodium chloride 84 per cent, potassium nitrate 83 per cent, silver nitrate 81 per cent, and potassium chlorate 83 per cent. If one of the ions is univalent and the other bivalent the dissociation is less. Examples of salts of this class are potassium sulphate 72 per cent, sodium sulphate 70 per cent, disodium phosphate 73 per cent, barium chloride 77 per cent, and zinc chloride 73 per cent. If both ions are bivalent the ionization is still less; copper sulphate 39 per cent, zinc sulphate 40 per cent. 600. The Hydrolysis of Salts. The fact that the salts of weak acids or bases are hydrolyzed by water has been frequently men- tioned. It will be recalled that the hydrolysis of salts is the oppo- site of neutralization. It is advisable at this point to review reactions of this kind in the light of what has been recently learned. Neutralization of an acid by a base is brought about as the result of the union of hydrogen and hydroxyl ions. Water itself 512 INORGANIC CHEMISTRY FOR COLLEGES is ionized to the extent of 0.00001 per cent, and an equilibrium exists between the molecules and the ions produced: 0.00001% H 2 O ;= H + + OH~ When hydrogen and hydroxyl ions are brought together by adding a solution of a base to one of an acid, the ions react to form water until this equilibrium is established. The reaction between sodium hydroxide and hydrochloric acid can be formulated as follows : Na + + OH- + H + + OP Na + + H 2 + Cr Since sodium and chlorine ions do not take part in the reaction if the solutions are sufficiently dilute, the equilibrium is between H + and OH~ on one side and H 2 O on the other, and the product of the concentration of these ions is the same as in pure water. When the salt of a weak acid is dissolved in water, hydrolysis takes place, that is, a part of the salt is converted into acid and base by the water. This can be illustrated by a consideration of the hydrolysis of sodium hypochlorite. Hypochlorous acid is dissociated in one-tenth normal solution at 18 to the extent of 0.02 per cent. There is, accordingly, an equilibrium established between the acid and its ions: 0.02% HCIO ^ H + + cicr If C10~ ions and H + ions are present in the same solution this equilibrium must be established. This occurs when sodium hypo- chlorite is dissolved in water, for the former furnishes C1O~~ ions and the latter H + ions. As a consequence, under these conditions three equilibria are set up which may be represented as follows: H 2 ^ H + + OH~ + NaOCl ^ C1CT + Na + 11 HCIO Some hydrogen ions from the water unite with the C10~ ions to form HCIO, and, as a consequence, an equivalent quantity of OH~ ions are left in the solution, which shows, therefore, an alka- THE PROPERTIES OF OXIDES, HYDROXIDES, AND SALTS 513 line reaction. The complete reaction may be represented as follows : Na + + CIO" + H 2 O ^ Na + + HC1O + OH~ Water reacts with the salt to form the base and acid, but as the latter is only slightly dissociated free hydroxyl ions are present in the solution which is, accordingly, alkaline. To sum up the subject to this point briefly, neutralization results from the union of H + + OH~ to form undissociated water, and hydrolysis results from the union of H + and an acid radical to form molecules of the undissociated acid. The extent to which a salt of a weak acid is hydrolyzed is deter- mined by the strength of the acid; the weaker the acid the greater the hydrolysis, because more hydrogen ions will be required to convert the negative ion into the undissociated acid, and more hydroxyl ions will be left free in the solution. A similar explanation can be given of the hydrolysis of salts derived from weak bases: in this case the undissociated base is formed and free hydrogen ions are left in solution. In the case of aluminium sulphate, for example, the reaction is as follows: A1 2 (SO 4 )3 + 6H 2 O ^ 2A1(OH) 3 + 3H 2 SO 4 If the salt is one of a weak acid and a weak base the extent of the hydrolysis is greater because there are two active agencies which bring it about the formation of both the undissociated base and the undissociated acid. The degree of hydrolysis is affected not only by the strength of the acid and base involved, but also by the concentration of the solution the more water the greater the hydrolysis and by the temperature. For example, copper sulphate is hydrolyzed suffi- ciently at room-temperature to show an acid reaction with litmus paper, but the hydroxide of the metal is not formed in large enough quantity to precipitate. If a dilute solution of the salt is boiled, more extensive hydrolysis takes place and a basic salt formed as the result of the partial hydrolysis of the neutral salt is slowly precipitated. A solution of magnesium chloride at 100 is sufficiently hydrolyzed to produce a concentration of hydrochloric acid which will attack iron; it is for this reason that water to be used in boilers should be freed from this salt. 514 INORGANIC CHEMISTRY FOR COLLEGES The relation between the degree of hydrolysis and the strength of acids and bases is indicated by the following examples: Borax is hydrolyzed to the extent of about 0.5 per cent in one-tenth normal solution; the ionization of boric acid at this concentration is 0.01 per cent. Aluminium sulphate in 0.001 N solution is hydrolyzed to the extent of 4.5 per cent at 25. Aluminium car- bonate and sulphide are both completely hydrolyzed by water. EXERCISES 1. Draw up a general statement of the physical and chemical properties of the hydroxides of metals having the valence 1, 2, and 3. 2. If the hydroxide of a metal has amphoteric properties would you expect it to be a strong base or a strong acid? 3. If you were given a sample of a chemical element how could you deter- mine by (a) physical and (6) chemical means whether it was metallic or non-metallic in character? 4. Give examples of the effect of change in valence of (a) acid-forming elements and (6) base-forming elements on the chemical properties of their compounds containing oxygen and hydrogen. 5. Would you expect that the chloride of an acid-forming element would be hydrolyzed by water? Give several examples. 6. From the facts given in this chapter state how you could distinguish the following from each other: (a) Fe(OH) 2 and Fe(OH) 3 , (6) Ca(OH) 2 , and Zn(OH) 2 (c) A1 2 (SO 4 ) 3 and ZnSO 4 , (d) Zn(OH) 2 and Mn(OH) 2 , (e) Cu(OH) 2 and Fe(OH),, (/) AgCl and PbCl,, (g) Na 2 CO 3 and NaCl. 7. Show by chemical equations, using ionic symbols, what happens when ammonium hydroxide is added to a solution of calcium phosphate in nitric acid. 8. (a) How do you think the two oxides of tin, SnO and SnO 2 , would differ in chemical properties? (6) Make a statement as to the behavior of SnCl 2 and SnCl 4 with water. 9. Compare the properties of the compounds which contain (a) anti- mony with the valence 3 and (6) with the valence 5. 10. When Cr 2 Os is heated in the air it is stable. If it is mixed with solid sodium hydroxide and heated in the air it is converted into sodium chromate, Na 2 CrO 4 , in which chromium has the valence 6. Can you explain why the reaction takes place? CHAPTER XXXV SODIUM, POTASSIUM, RUBIDIUM, AND CAESIUM 601. Compounds of sodium and potassium are widely dis- tributed, play an important part in natural processes, and are used extensively in the arts. Lithium, rubidium, and caesium, the other members of the first family of the first group in the periodic classification of the elements, occur in but small quantities in nature and the uses to which they have been put are very limited. The elements of this family are the most active of the metals; they oxidize quickly in the air, and are, therefore, kept under oil; they decompose water rapidly and form soluble hydroxides, which are strong bases and are called caustic alkalies. Their salts are very stable and resist high temperatures. The hydroxide of caesium is the strongest base, a fact in accord with the conclusion already stated that, in most cases, increase in atomic weight in a chem- ical family is associated with increase in basic properties. The relationships between the chemical and physical properties of potassium, rubidium, and caesium and their compounds are very close, and there is a gradation in these properties similar to that emphasized in the case of the halogens. Lithium and sodium, however, do not fit so well into the family from this point of view; the divergence here is similar to that noted in the case of fluorine and the other members of the halogen family, and nitrogen and the phosphorus family (449). The melting- points of the elements decrease from lithium, 186, to caesium, 26.5, but the specific gravity of sodium, 0.97, is greater than that of both lithium, 0.53, and potassium. 0.86. The chlorides of potassium, rubidium, and caesium are isomorphous with one another, but not with sodium chloride. Only the more impor- tant compounds of sodium and potassium will be considered below. 515 516 INORGANIC CHEMISTRY FOR COLLEGES SODIUM 602. Sodium occurs in sodium chloride, complex silicates, borax, Chile saltpeter, cryolite, and in other minerals. It is present in sea-plants as salts of organic acids, which are con- verted into sodium carbonate when the plants are burned. Sodium chloride is an important constituent of the blood. Sodium was first isolated in 1807 by Davy, who obtained it by the action of an electric current on moist sodium hydroxide. The metal is manufactured by a process invented by Castner. Fused sodium hydroxide is electrolyzed in an iron vessel in which is suspended the anode, a, (Fig. 37). Over the cathode, c, is placed a drum, d, the lower part of which is made of iron gauze, g. Oxygen is evolved at the anode and rising through the fused mass escapes at e. Sodium and hydrogen liberated at the cathode collecting; the gas and the molten metal are led off through pipes. As the electrolysis proceeds fresh alkali is added through the hole at e. When the cell is started the hydrox- ide is fused by gas jets at i, but later the passage of the current furnishes the heat required to keep the electrolyte molten. The chief physical and chemical properties of sodium have already been given (page 443) . It is so soft it can be readily cut with a knife. It dissolves in liquid ammonia, and in mercury (671). It unites with the non-metals, and forms a crystalline compound with hydrogen at 365, sodium hydride, NaH, which reacts with water to form sodium hydroxide and hydrogen. Sodium is used in the preparation of sodium peroxide, sodium cyanide, NaCN, and many organic compounds. It is also used in the laboratory for drying liquids, such as ether, which do not react with it. 603. Sodium Chloride. Common salt and Chile saltpeter are the sources from which all the compounds of sodium are made. FIG. 37. SODIUM, POTASSIUM, RUBIDIUM, AND CAESIUM 517 The nitrate is used in making nitric acid and the sodium sulphate formed as a by-product is used as such or in making glass and other substances. Salt, on account of its cheapness, is, however, the chief source of sodium compounds. It is obtained as rock salt and from sea-water, or from salt-brines derived from lakes or wells. Rock salt is found, at times, in a very pure condition and, after being mined, is ground and used without purification. It occurs in the United States, England, Austria, Germany, and Spain in large deposits. Rock salt is usually contaminated with oxides of iron, clay, sand, and other substances, and is purified by crystallization. It is mined in New York, California, Kansas, Utah, and Louisiana. The chief sources of salt in the United States are brines found in New York near Syracuse and Warsaw, in Michigan at Saginaw Bay and Manistee, and near Salina, Kansas. The brine occurs some distance below the surface and is obtained by pumping from wells, which are bored about 8 inches in diameter and lined with iron casings. The brine is treated with milk of lime or sodium carbonate to precipitate calcium and magnesium compounds; it is then sepa- rated and evaporated to obtain the salt. In the commercial process complete purification is not effected, and as a consequence, ordinary salt is more or less deliquescent. Sodium bicarbonate is sometimes added to table-salt to convert the magnesium chloride present into the insoluble carbonate, which does not absorb water from the air. Very pure sodium chloride may be prepared for chemical purposes by conducting hydrogen chloride into a satu- rated solution of salt until no more of the gas dissolves. Salt is practically insoluble in concentrated hydrochloric acid, and is precipitated, while the impurities remain in solution. Salt crystallizes in cubes, the faces of which are usually hollow. As the faces grow, layer by layer, they include mechanically between them a small amount of the liquid from which they are formed. As a consequence, when the crystals are heated to a high temperature the liquid is vaporized and the steam produced causes them to fly apart with a crackling noise; the salt decrepitates. Salt melts at 20 and boils at 1750. Its solubility is given in the diagram on page 506. Salt is a necessary constituent of the diet of animals, as it plays an important part in the blood; it is used as the source of chlorine and hydrochloric acid, in the man- 518 INORGANIC CHEMISTRY FOR COLLEGES ufacture of sodium compounds, in the preservation of meat and fish, in glazing pottery, and for other purposes. 604. Sodium Hydroxide. This compound, which is commonly called caustic soda in trade, is manufactured by the electrolysis of sodium chloride or by treating a solution of sodium carbonate (" soda-ash ") with milk of lime, which is a suspension in water of finely divided calcium hydroxide: Na 2 CO 3 + Ca(OH) 2 = 2NaOH + CaCO 3 Steam is blown into the mixture in order to agitate the solid and thus keep it in contact with the solution. When the reaction is complete, the calcium carbonate is allowed to settle, and the dilute solution of sodium hydroxide drawn off and evaporated, first in vacuum kettles until the impurities crystallize out, and then in iron pots over an open fire. When all the water has been driven off, the fused sodium hydroxide is poured into iron drums. Large quantities of caustic soda are manufactured by the elec- trolysis of a solution of sodium chloride. The process involves the use of electrical energy, but the cost of this is offset, in part, by the fact that sodium chloride is used rather than sodium car- bonate; the latter must first be manufactured from salt and is, therefore, the more expensive substance; and, further, the chlorine formed at the same time is a valuable by-product. 605. When an electric current is passed through a solution of sodium chloride, sodium hydroxide and hydrogen are produced at the cathode and chlorine at the anode. If the two portions of the solution are allowed to mix, the halogen and alkali react and sodium hypochlorite is formed. The large number of cells which have been invented to prepare sodium hydroxide and chlorine electrolytically differ from one another chiefly in the way in which the formation of the hypochlorite is avoided. In one type of cell a diaphragm is used to separate the liquid around the cathode from that around the anode. The diaphragm is constructed of porous material, which allows the liquid to penetrate it, and thus does not prevent the flow of the electric current; but it does pre- vent mechanical mixing of the solutions in the two compartments. One form of diaphragm used is made of a mixture of asbestos and iron oxide supported on an iron grating, which serves as the cathode. The Townsend cell, which is an example of this type, is represented SODIUM, POTASSIUM, RUBIDIUM, AND CAESIUM 519 sHydrogen '* Outlet MiL%5: LJ -JBfmeJ\ Lye Outlet i 'Connection to Cathode diagrammatically in Fig. 38. A is the anode, which is made of graphite. A solution of a chloride is allowed to flow slowly into the cell, and is kept at the level indicated. The cathode is of perforated iron and supports the diaphragm. Kerosene is placed in the outer compartment. /+ When the electrolysis is taking nn ..-Anode place the solution slowly per- colates through the diaphragm due to the hydrostatic pressure of the water; it falls in drops through the kerosene to the V bottom of the outer compart- ment from which it passes out through a pipe. The solution obtained contains about 150 grams of sodium hydroxide and 200 grams of sodium chloride FIG. 38. per liter. It is evaporated until the salt crystallizes out, separated from the solid, and finally heated in iron pots until all the water has been driven off. The molten sodium hydroxide is then poured into iron containers. The Allen-Moore and Nelson cells resemble in construction the Townsend cell, but kerosene is not used in the outer com- partments. 606. The Castner-Kellner cell (Fig. 39) is an example of a different type, in which no diaphragm is used. The anodes are FIG. 39. of graphite, and are placed in compartments separated by par- titions from the compartment containing the cathode, which is made of iron. A layer of mercury is placed on the bottom of 520 INORGANIC CHEMISTRY FOR COLLEGES the cell as indicated. The chlorine liberated escapes through the outlet indicated in the diagram. Metallic sodium is set free at the surface of the mercury in the anode compartments and dissolves in the mercury. During the electrolysis the cell is rocked slowly by an eccentric, and the mercury flows from the anode to the cathode compartment and back again. The depres- sions at the bottom of the tank are of such a size that during the rocking none of the solution from the anode compartment gets into the cathode compartment. When the mercury containing the sodium in solution reaches the cathode compartment, the sodium reacts with water to form sodium hydroxide, and hydrogen is evolved at the cathode. A strong solution of sodium hydroxide free from salt is obtained. 607. Sodium hydroxide is a white solid which absorbs water from the air, and serves, therefore, as a valuable drying agent. It melts at red-heat. It is used in large quantities in the man- ufacture of soap, paper-pulp, and phenol, and other important organic compounds used in the preparation of dyes. 608. Sodium Carbonate. The mineral called trona, which has the composition NaHCO3,Na2CO3,2H2O, and the ashes from sea- plants were used as the sources of sodium carbonate until the development of the chemical industries made it imperative to invent a process for the preparation of this very important com- pound from salt. As the result of war France was cut off from a supply of the carbonate and a prize of 100,000 francs was offered by the French Academy for the best solution of the problem. Le Blanc invented a process, which has since been known by his name, won the prize, and was granted a patent in 1791. During the French Revolution his factory was seized and declared to be public property. He was paid no indemnity and, finally, in discouragement, committed suicide. The process, with but slight improvement, is used extensively to-day, although in the United States it has been superseded by another process invented by Solvay, a Belgian chemical manu- facturer. In both cases sodium chloride is the source of the sodium, and calcium carbonate is the source of the carbonate radical, but the reactions used to effect the formation of sodium carbonate are different. The value of the by-products produced largely determines the net cost of the carbonate and, consequently, SODIUM, POTASSIUM, RUBIDIUM, AND CAESIUM 521 the process to be used under the conditions which exist in the chemical market. It will be of interest, therefore, to consider in some detail these two processes to see how the costs of materials and energy and the value of by-products are the most important factors in industrial chemistry from the economic point of view. 609. The Le Blanc Process. It will be recalled that when a solution of sodium carbonate is added to one of calcium chloride a double decomposition takes place owing to the insolubility of calcium carbonate; Na 2 CO 3 + CaCl 2 = CaC0 3 + 2NaCl Since sodium chloride does not interact with calcium carbonate, the reaction cannot be made to proceed in the reverse direction. Le Blanc's solution of the problem was to convert the chloride into another salt, which, when fused with calcium carbonate, entered into double decomposition with it. Sodium sulphide acts in this way: Na 2 S + CaCOs = CaS + Na 2 CO 3 The sulphide was prepared by converting the chloride into the sulphate, 2NaCl + H 2 SO 4 = Na 2 SO 4 -f 2HC1, and fusing the latter with carbon: Na 2 SO 4 + 2C = Na 2 S + 2CO 2 The products used are, accordingly, salt, coal, calcium carbonate, and sulphuric acid; the by-products are hydrochloric acid, cal- cium sulphide, and carbon dioxide. The economic success of such a process depends upon the utilization of these by-products; sul- phuric acid must be manufactured and is, thus, an important item in the expense. At first the hydrochloric acid formed was a nui- sance, for it was discharged into the air, but as the chemical indus- tries developed it found important uses and became, therefore, of value. The acid not required for these purposes is converted into chlorine, which is largely used to make bleaching powder. In this way this by-product becomes of value, and serves to reduce the cost of the sodium carbonate. 522 INORGANIC CHEMISTRY FOR COLLEGES The sulphur in the calcium sulphide is recovered by what is known as the Chance-Glaus process. The sulphide is suspended in water and treated with carbon dioxide, which sets free hydrogen sulphide : CaS + H 2 + C0 2 = CaC0 3 + H 2 S The latter is then burned in a furnace with just enough air to con- vert it into sulphur: 2H 2 S + O 2 = 2H 2 + 2S The sulphur is sold or made into sulphuric acid to be used in the process. About 85 per cent of the sulphur in the sulphuric acid used in converting the salt into sodium sulphide can be recovered by this process. 610. The Solvay Process. The reactions upon which this proc- ess is based were discovered in 1838. Attempts were made early to utilize the reactions for industrial purposes, but they were unsuccessful, largely because the process involved the handling of gases on the large scale and the careful regulation of temperature, the technique of which had not been fully worked out. Solvay proved himself to be a capable engineer and overcame the difficul- ties. In 1863 he designed the necessary apparatus for a plant and in 1873 the process had established itself as a commercial success. It is the one used in the United States for the manu- facture of sodium carbonate. In this process sodium chloride is first converted into sodium bicarbonate, which is difficultly soluble in water, by treating it with ammonium bicarbonate, which is soluble: NaCl + NH 4 HCO 3 = NaHCO 3 + NH 4 C1 When this salt is heated it decomposes: 2NaHCO 3 = Na 2 CO 3 + H 2 O + CO 2 The ammonium bicarbonate required is made by the action of car- bon dioxide on a solution of ammonia in water: NH 4 OH + H 2 C0 3 = NH 4 HC0 3 + H 2 The carbon dioxide is obtained by heating limestone, CaCOs = CaO + CO 2 ; and the lime formed is used to recover the ammonia SODIUM, POTASSIUM, RUBIDIUM, AND CAESIUM 523 from the ammonium chloride produced according to the first reaction given above: 2NH 4 C1 + Ca(OH) 2 = 2NH 3 + 2H 2 + CaCl 2 The products used in the process are, accordingly, sodium chloride and calcium carbonate, and the by-product is calcium chloride; the ammonia is recovered and used over and over again. It is seen, thus, that through the ingenious use of the fact that sodium bicarbonate is difficultly soluble in water, it is possible to realize the transformation indicated by the equation CaCO 3 + 2NaCl = Na 2 C0 3 + CaCl 2 which is the reverse of that which expresses the facts when sodium carbonate and calcium chloride interact in solution. The by-product of the Solvay process, calcium chloride, is of little value; it is cheap enough to warrant its use on roads to lay dust. This method of preparing sodium carbonate does not yield as by-products hydrochloric acid and chlorine as the Le Blanc process does. Chlorine is made in the United States by the electrolysis of salt, because water-power is available and electrical energy correspondingly cheap. It is this factor which determines the use of the Solvay process in America and that of Le Blanc in England. 611. Properties of Sodium Carbonate and Sodium Bicarbonate. Sodium carbonate crystallizes from water below 35.2 as a decahydrate, Na 2 CO 3 ,10H 2 O, which is sometimes called washing soda or sal soda. It effloresces rapidly in the air and is converted into a white powder which has the composition of the crystals formed from solutions above 35.2, namely, Na2CO 3 ,H 2 O. The anhydrous salt is called in trade soda-ash. Sodium carbonate is hydrolyzed in aqueous solution, 3.2 per cent being converted into the hydroxide in N/25 solution at 25: Na 2 CO 3 + H 2 O ;= NaHCO 3 + NaOH Many of the uses of sodium carbonate are based on this fact. It is used in washing and for scouring purposes because the small amount of alkali present has a marked effect on oil or grease; it converts them into the colloidal condition as the result of the 524 INORGANIC CHEMISTRY FOR COLLEGES formation of minute globules which become suspended in the solution and are, as a result, easily removed. Sodium carbonate is used in large quantities in softening water (629), in the manufacture of glass and soap, and in smaller amounts in many other industries. Sodium bicarbonate, NaHCOs, is an anhydrous salt; when a solution of the salt in water is boiled, it is hydrolyzed and carbon dioxide slowly escapes; sodium carbonate is formed to some extent and owing to the hydrolysis of the latter the solution becomes alkaline. Sodium bicarbonate has long been used under the name baking soda or saleratus. In making leavened bread it is necessary to add to the dough something that produces bub- bles of gas which expand when the dough is baked, and thus give to the bread a porous structure. Carbon dioxide is the gas com- monly used for this purpose; it is generated either through the use of yeast, which is a plant that evolves carbon dioxide as it grows, or as the result of the interaction of sodium bicarbonate and an acid put into the dough. Molasses and vinegar contain acids and can be used with the salt, or cream of tartar, acid potassium tartrate, KH(C4H 4 O6), may be employed. 612. Baking powders are mixtures of sodium bicarbonate and a substance which reacts with the latter to form carbon dioxide. The tartrate powders contain tartaric acid, H 2 (C 4 H 4 O6), or cream of tartar, the phosphate powders contain primary calcium or sodium phosphate, and the alum powders contain aluminium sulphate or a double salt of the latter and sodium sulphate. The reaction which takes place in the case of cream of tartar powders is as follows: NaHC0 3 + KH(C 4 H 4 O 6 ) - KNa(C 4 H 4 O 6 ) + H 2 O + CO 2 When aluminium sulphate is used, the salt first hydrolyzes to form aluminium hydroxide and sulphuric acid: A1 2 (SO 4 ) 3 + 6H 2 O = 2A1(OH) 3 + 3H 2 SO 4 and the latter decomposes the bicarbonate: 2NaHCO 3 + H 2 SO 4 = Na 2 SO 4 + 2H 2 O + 2CO 2 The products left in the bread are, accordingly, either potassium sodium tartrate, which is commonly called rochelle salt, or alumin- SODIUM, POTASSIUM, RUBIDIUM, AND CAESIUM 525 ium hydroxide, or, in the case of the phosphate powders, second- ary salts of the acid. Starch is added to baking powder to cover the grains of the chemicals used so that they cannot interact until the powder is moistened in the dough. It is evident that baking powders should be protected from the air which contains moisture. The starch also serves as a diluent, and the amount added to each kind of powder is that required to give a mixture which evolves about 12 per cent of its weight of carbon dioxide. As a result, the same amount of different powders have about the same leavening effect. Sodium bicarbonate is used in Seidlitz powders, bromo-seltzer, and other effervescing medicinal drinks. 613. Other Salts of Sodium. A number of sodium salts have already been described more or less fully. Among these are the following: Borax, sodium nitrite, sodium nitrate, sodium sul- phate, sodium thiosulphate, sodium sulphite, sodium silicate, sodium peroxide, and the sodium phosphates. 614. The Test for Sodium Salts. Two difficultly soluble salts are used at times in qualitative tests for the element; these are sodium fluosilicate, Na^SiFe, which is formed by adding hydro- fluosilicic acid to a strong solution of a sodium salt, and sodium- hydrogen pyroantimonate, Na2H2Sb2C>7, which is formed when the potassium salt of the acid is added to the solution. The test commonly applied is the examination of the spectrum of the light produced when some of the solution suspended on a platinum wire is put into a Bunsen flame. Sodium compounds give a bright yellow flame. If the characteristic lines persist for some time and are not, therefore, produced as the result of traces of sodium com- pounds as accidental impurities, the presence of the element is shown. The compounds of several elements yield colored flames when they are placed in a Bunsen flame. In order to determine what lements are present, the light is examined with a spectro- scope. In the quantitative determination of sodium, all other elements are separated and the metal is finally converted into sodium sulphate, which is heated to a high temperature and weighed. 615. Spectroscope. When a beam of white light passes through a triangular prism made of glass, it is broken up into its constituents, and a spectrum is obtained which shows the colors of 526 INORGANIC CHEMISTRY FOR COLLEGES the rainbow. This results from the fact that white light is made up of vibrations of different wave-lengths, each of which produces a characteristic color. If the light produced when a sodium salt is heated in a Bunsen flame is passed through a slit to produce a beam and then through the prism, a bright yellow line is obtained and the rest of the spectrum is lacking. If a potassium salt is used a number of blue lines are visible along with a few red and green lines. Each element in the free condition or in combination, when it is heated to a sufficiently high temperature, gives in a similar way a characteristic spectrum made up of colored lines. The spectra of the elements have been determined, and the presence of any one of the latter in a mixture can be dis- covered by an examination of the spectrum produced when the material containing it is heated to such a temperature that the characteristic vibrations are produced. The optical instrument by means of which the light is examined is called a spectroscope. It consists essentially of a tube furnished at one end with a slit through which the light passes, and at the other end of which is a lens to convert the light into a parallel beam. The latter then passes through a triangular prism and the issuing beam, which is spread out into its constituents, is examined by means of a small telescope. In the direct vision spectroscope the lenses and prisms are so arranged that but one tube is necessary. At the temperature produced in a Bunsen flame only the spectra of the metals of the alkalies or alkaline earths can be observed. The chlorides of the metals are commonly used on account of their volatility. POTASSIUM 616. The occurrence of potassium in feldspar and other complex silicates has already been noted (556). The chief sources of the element and its compounds are the chloride and sulphate present in the salt deposits in Alsace and Stassfurt, Germany, where they occur as sylvite, KC1, carnalite, KCl,MgCl2,6H2O, and kainite, K2SO4,MgSO4,MgCl2,6H2O. Smaller quantities are obtained from wood ashes, the residue from beet sugar, and wool washings. Largely as the result of the recent war the enormous beds of sea- weed along the coast of California have been used as a source of potassium and of iodine. The ashes obtained by burning the sea- SODIUM, POTASSIUM, RUBIDIUM, AND CAESIUM 527 weed contain about 9 per cent potassium chloride and 0.1 per cent iodine as potassium iodide. Large deposits of alunite, K 2 SO4,A1 2 (S04)3,4A1(OH)3, exist in Utah, Nevada, and Colorado, and are now utilized as a source of potassium sulphate. When the mineral is roasted the sulphate and hydroxide of aluminium are decomposed into oxide, and the product on treatment with water yields a solution of potassium sulphate. Potassium chloride has recently been recovered successfully from the fumes from cement plants and blast furnaces. Although much attention has been paid as a result of the war to the production of potassium com- pounds from the natural resources of the United States, it is highly probable that Europe will for some time supply the world with potash salts, which are essential plant-foods, and, therefore, important constituents of fertilizers. Potassium was first isolated by Davy in 1807; it is prepared by a method similar to that used in the case of sodium. Its physical and chemical properties resemble those of sodium (602). 617. Potassium hydroxide is made by the electrolysis of potas- sium chloride, and resembles closely sodium hydroxide. When it is required in a very pure form it is dissolved in alcohol in which the chloride and carbonate, the common impurities present, are insoluble. After the clear liquid has been separated from the in- soluble material, it is evaporated to dry ness. Potassium hydroxide is very deliquescent and very soluble in water, 1 liter of which will dissolve about 1500 grams of the solid. It is called caustic potash in trade. It, like sodium hydroxide, decomposes proteins and dis- solves the flesh if left in contact with it. 618. Salts of Potassium. Potassium carbonate cannot be made from potassium chloride by the Solvay process because potassium bicarbonate is readily soluble in water. It is for this reason that a solution of potassium hydroxide is preferred to one of sodium hydroxide as an absorbent for carbon dioxide. When the sodium hydroxide is used, the bicarbonate separates after a time and clogs up the tubes, etc., used in the absorbing apparatus. The carbonate is made by the Le Blanc process or in other ways. Its trade name is potash or pearlash. It is prepared in appre- ciable quantities from suint, the fatty material present in wool, which contains along with grease the potassium salt of an organic acid. The grease is extracted and purified and sold under the 528 INORGANIC CHEMISTRY FOR COLLEGES name " lanolin," and the residue is ignited and converted into potassium carbonate. The salt is usually sold in the anhydrous form, and differs from sodium carbonate in being very deliquescent. The crystals obtained from water have the composition 2K2COs,- 3H2O. It is used in the manufacture of glass, soft soap, and certain potassium salts. 619. Potassium nitrate was formerly prepared from the products resulting from the decomposition of animal refuse in natural or artificial nitrate beds, but during the Crimean war (1852-55) the supply of this important ingredient of gunpowder became deficient and the method now used was invented. This involves the treat- ment of potassium chloride with sodium nitrate in aqueous solu- tions : KC1 + NaNO 3 = KNO 3 + NaCl The reaction has already been explained (592). Potassium nitrate (saltpeter or niter) melts at 339 and at higher temperatures is reduced to potassium nitrite. Potassium nitrate is used in the manufacture of gunpowder, matches, and fireworks, the use depending on the fact that the salt gives off oxygen freely at high temperatures and is, therefore, an active oxidizing agent. It is also used in curing meats; it prevents putrefaction and produces the deep red color familiar in the case of salted hams and corned beef. 620. Gunpowder is the oldest explosive and is still a very important one. According to tradition it was invented by a monk at Freiburg, Germany, in the fourteenth century; it was first used at the battle of Crecy in 1346. Gunpowder is a mixture of potas- sium nitrate, sulphur, and charcoal. Its explosive properties are due to the fact that when it is ignited there is produced a large volume of gas, which consists principally of nitrogen and carbon dioxide. A mixture of potassium nitrate and charcoal burns vigorously when ignited, but a large part of the carbon unites with the potassium to form potassium carbonate. In order to prevent this and thus increase the proportion of gas formed, sulphur is added to the mixture in sufficient quantity to unite with the metal to form potassium sulphide. When gunpowder burns in the open air a reaction takes place which is indicated by the following equation : 2KNO 3 + 3C + S = K 2 S + 3CO 2 + N 2 SODIUM, POTASSIUM, RUBIDIUM, AND CAESIUM 529 When it is exploded in a confined space the reaction is much more complex, other substances in addition to those just mentioned being formed, among which are potassium carbonate, potassium sulphate, and carbon monoxide. The products consist of about 43 per cent gases and 57 per cent solids by weight. 621. Other Salts of Potassium. A large number of potassium salts are known, and many are used in the laboratory in preference to the cheaper sodium salts when the acid radical of the salt is desired. For example, the iodide commonly used is potassium and not sodium iodide. The more expensive salts are employed because in most cases they are less soluble in water, and, therefore, more easily purified by crystallization, and because they are more stable in the air. Certain sodium salts are hygroscopic whereas the corresponding potassium salts are not; in the case of the carbonates, however, the potassium salt is hygroscopic and, as a consequence, sodium carbonate is the ordinary laboratory reagent. Potassium cyanide, KCN, is made by heating potassium ferro- cyanide (751). It yields hydrocyanic acid (prussic acid), b.p. 20.5, when treated with an acid. The salt and the acid are active poisons. 622. Tests for Potassium Salts. The presence of potassium salts is tested for qualitatively by observing the flame produced by introducing into a Bunsen flame some of the salt on a platinum wire. If a spectroscope is not available, the flame is examined by looking at it through a piece of blue cobalt-glass, which allows the violet light produced by potassium to pass through it but cuts off the light from sodium and other metals. Potassium is separated from sodium quantitatively by adding to the solution of its salts chloroplatinic acid, H^PtCle, which precipitates potassium chloroplatinate, K^PtCle. Since the salt is somewhat soluble in cold water, it is precipitated and washed in a mixture of water and alcohol in which it is less soluble. Potassium perchlorate is very slightly soluble in water and is also used in the quantitative determination of potassium. EXERCISES 1. Starting with sodium chloride as the source of sodium, indicate by chemical equations by what reactions the following compounds can be pre- pared: (a) NaNO,,"(6) Na 2 B 4 O 7 , (c) Na 2 HPO 4 , (d) Na 2 S 2 O 7 , (e) Na 2 O 2 , (/) N a2 S 2 3 . 530 INORGANIC CHEMISTRY FOR COLLEGES 2. (a) Why can NaCl be readily removed from many other salts by crystallization? How would you proceed to get (6) pure NaCl, and (c) KNO 3 , from a mixture of the two salts? 3. State the economic advantages of combining a plant for making ammonia by the HaBer process with one making sodium carbonate by the Solvay process provided NH^Cl is used instead of (NH 4 ) 2 SO4 in fertilizers. Consider raw materials and by-products. 4. Why cannot Na 2 CO 3 replace NaHCO 3 in baking powders? 5. Calculate (a) the weight of cream of tartar required to react with 100 grams NaHCO 3 when the two substances are mixed to prepare a baking powder. Calculate (6) the weight of CO 2 given off by the mixture (c) the percentage by weight of CO 2 given off by the mixture, and (d) the weight of starch that must be added so that the resulting mixture of the three ingredi- ents will yield 12 per cent of its weight of CC>2. 6. Starting with alunite, indicate by equations by what chemical reac- tions the following could be prepared: (a) K 2 SO 4 , (6) A1C1 3 , (c) KC1, (d) alum (K 2 S04, A1 2 (SO 4 ) 3 , 24H 2 0). CHAPTER XXXVI CALCIUM, STRONTIUM, BARIUM, AND RADIUM 623. Calcium, strontium, and barium form a typical family of elements which are known as the alkaline earth metals; their atomic weights are, respectively, 40.07, 87.63, and 137.37. They always have the valence 2, and the general chemical behavior of their compounds is that shown by the compounds of other metals when they exhibit this valence. We shall see as the chemistry of the metallic elements is developed, that the valence shown by a metal is a most important factor in determining the behavior of its salts. Barium, the member with the highest atomic weight in the family, is the most active in base-forming properties. Like potas- sium and sodium it forms a peroxide, BaO 2 , when heated in the air; its hydroxide is decomposed only at high temperatures into oxide and water; its nitrate is reduced by heat to a nitrite, which, unlike the nitrites of potassium and sodium, can be converted at a high temperature into the oxide and oxides of nitrogen. The sulphate and carbonate of barium require high temperatures to convert them into oxides. In the stability of its compounds toward heat, barium stands below potassium and sodium, but resembles them closely. The compounds of strontium and calcium behave in a similar way, but lower temperatures are required to effect the changes, which take place more readily in the case of calcium. The solubilities of the corresponding salts of the three ele- ments change progressively. In the case of the hydroxides the solubility decreases in the order barium, strontium, calcium, the weight of the base dissolving in 100 grams of water at 18 being, respectively, 3.7, 0.77, and 0.17grams. The values of the solubilities of the sulphates are progressive, but in the reverse order; they are, beginning with barium sulphate, 0.0s23, 0.011, and 0.20 gram in 100 grams of water at 18. The carbonates of the three elements 531 532 INORGANIC CHEMISTRY FOR COLLEGES are exceedingly insoluble and differ but little from one another in this respect; the solubility of calcium carbonate is given as 0.0s 13 gram in 100 grams of water. CALCIUM 624. The more important forms in which calcium occurs in nature have been stated (556) ; over 3 per cent of the earth's crust is calcium, which stands fifth in the list of elements when they are arranged according to their abundance (556) . Calcium was first isolated in an impure condition by Davy in 1808 and was prepared in 1898 by Moissan by heating calcium oxide with sodium. It is manufactured now by the electrolysis of fused calcium chloride. The more important physical properties of calcium are given in the table on page 443. It is an excellent conductor of elec- tricity. When heated it unites with hydrogen to form calcium hydride, CaEb, and with nitrogen to form a nitride, Ca3N2, which decomposes with water to form ammonia and calcium hydroxide. The metal, which has become only recently available commercially, has not yet found any industrial uses. 625. Calcium Chloride. Sea-water contains calcium chloride and, consequently, the compound is found in the salt deposits formed as the result of the evaporation of inland seas. It occurs in combination with magnesium chloride in the mineral tachydrite, CaCl 2 ,2MgCl 2 ,12H 2 O, at Stassfurt. Calcium chloride crystallizes as a hexalrydrate, CaCl2, 6H2O, from water. The salt is very deliquescent, and on account of its great solubility it is used in freezing mixtures. When heated to a sufficiently high temperature it loses its water of crystallization, but is at the same time partially hydrolyzed and the resulting product contains some calcium hydroxide or calcium oxide. In preparing the desiccated material to be used as a drying agent the hexahydrate is heated not higher than 200 and converted into the dihydrate, CaCl2, 2H20, which is the so-called granulated cal- cium chloride used in the laboratory. Even this contains some hydroxide and if it is to be used in the quantitative determination of water when carbon dioxide is present, the salt is first exposed to a stream of dry carbon dioxide to convert the hydroxide into car- CALCIUM, STRONTIUM, BARIUM, AND RADIUM 533 bonate. Dehydrated calcium chloride unites with ammonia to form a salt, CaChjSNHs, with organic derivatives of water, such as the alcohols, and with the compounds derived from ammo- nia. For this reason it cannot be used to dry these substances. 626. Calcium Carbonate. Limestone, marble, coral, and chalk occur in vast quantities on the earth's surface; the mountains of Switzerland are composed largely of limestone and those of the Dolomite region in Central Europe of a mineral which is called dolomite, CaCOsjMgCOs, from the name of the scientist who first described the region fully. Limestone in the purest condition consists of minute crystals of calcium carbonate which are asso- ciated with varying amounts of iron oxide, silica, clay, magnesium carbonate, and other substances; its color may be white, gray, blue, or black depending upon the impurities present. Limestones have been produced from the remains of animal organisms which existed in the sea; they contain, at times, shell formations, which can be clearly seen. The harder varieties of limestone are exten- sively used as a building material, but they are attacked slowly by air and water and are not as permanent as sandstone or granite for this purpose. Marble is a more or less pure, highly crystalline, massive variety of calcium carbonate, which can be polished. Onyx marbles are translucent and are made up of bands, colored by impurities, which have been deposited from cave or spring waters. Marble is an excellent building stone; as it is much purer than many limestones it is not readily attacked by atmospheric influ- ences. The red, pink, or green color of certain marbles is largely due to iron oxide; the gray and black marbles are colored by carbonaceous material. Chalk and coral are made up of calcium carbonate derived from the calcareous parts of marine organisms. Calcium car- bonate is abundant in parts of England, and forms the well-known chalk cliffs along its southern coast. The shells of eggs, oysters, clams, and mussels are composed largely of calcium carbonate. Pearls are made up of the carbonate deposited in exceedingly thin layers in such a way that when light falls upon them it is broken up into the colors of the spectrum. Calcium carbonate occurs in well-defined individual crystals as calcite, Iceland spar, and aragonite. The latter belongs to the 534 INORGANIC CHEMISTRY FOR COLLEGES rhombic system, and is an unstable variety; when heated gently it changes to a mass of minute crystals of calcite. Iceland spar and calcite belong to the hexagonal system. When calcium car- bonate is precipitated it is formed in an amorphous condition, but on standing it becomes crystalline and less soluble in water. When substances are precipitated in making quantitative analyses, the precipitate is usually heated for some time in contact with water in order to convert it into the crystalline form in which it is less soluble; owing to its granular condition, the precipitate is more readily separated from the solution by filtration. Calcium carbonate is used in the manufacture of lime, glass, and sodium carbonate, as a flux in metallurgical operations (557), and is added to soils which have become acid in reaction as the result of the formation of organic acids from decaying vegetable matter. Calcium carbonate is a by-product of many chemical industries. It is sold under the name of " whiting," which is prepared by grinding and washing chalk, and is used to modify the shade of pigments, in polishing powders, and for other purposes; mixed with about 18 per cent of linseed oil it forms putty. Pre- cipitated calcium carbonate is the chief ingredient of tooth powders. 627. Calcium Bicarbonate. Calcium carbonate dissolves in water containing carbon dioxide as the result of the formation of an acid salt: CaC0 3 + H 2 C0 3 ^ Ca(HCO 3 ) 2 The reaction is a reversible one; the acid salt is formed at ordinary temperatures, but if the solution containing it is heated or evap- orated the carbonate is precipitated and the carbonic acid breaks down into carbon dioxide and water. The reaction is often observed in regions containing limestone or marble. Carbon dioxide is frequently formed within the earth and is found in the waters of effervescing springs. When such water comes into contact with masses of calcium carbonate the latter is slowly dissolved; some caves have been formed in this way. The beautiful formations in the caves found in limestone regions have been produced by the decomposition of the acid carbonate contained in the water which has dissolved calcium carbonate. As such water drops from the roof of the cave it slowly evaporates, and the deposits of calcium carbonate formed are called stalactites CALCIUM, STRONTIUM, BARIUM, AND RADIUM 535 if they hang from the roof, or stalagmites when they accumulate on the floor. 628. Hard Water. The action of carbon dioxide and water on magnesium carbonate and ferrous carbonate is similar to that just described in the case of the calcium salt. Water which contains salts of calcium, magnesium, or iron in solution is said to be " hard," because when it is used with soap the solution does not have the smooth feeling characteristic of soap solutions. When hard water of this type is boiled, the bicarbonates present are decomposed, the insoluble normal carbonates are precipitated, and the carbon dioxide escapes as a gas; the water is thus rendered "soft" and dissolves soap. Hard water which can be made soft by boiling is said to be " temporarily" hard. Calcium sulphate occurs as a mineral and as it is slightly soluble in water it is often found in water-supplies. In this case boiling the water does not remove the salt and for this reason the water is said to be " permanently" hard. Hard water of either type cannot be used with soap for cleansing purposes until the hardness is removed. When soap dissolves in water, the solution produced converts the grease present into an emulsion which consists of globules of microscopic size that stay sus- pended in the water. There is formed what appears to be a cloudy solution, which can be washed away by more water. If the water is hard, the salts in solution react with the soap and form with it insoluble compounds which are precipitated. As more soap is added the process continues until the metallic ions have been precipitated; the soap then stays in solution and produces its normal effect on the grease. 629. Common soap consists of a mixture of the sodium salts of certain organic acids which are obtained from fats (502); among these are stearic acid, which will serve as an example. The salts of this acid which contain calcium, magnesium, or iron are insoluble in water, and, as a consequence, when soap is added to hard water a double decomposition takes place, which is repre- sented in the case of sodium stearate as follows: CaSO 4 + 2Na(Ci 8 H35O 2 ) = Ca(Ci 8 H 3 5O 2 )2 + Na 2 SO 4 Calcium stearate is precipitated when the soap is added to water as long as the latter contains a calcium salt. 536 INORGANIC CHEMISTRY FOR COLLEGES The removal of the metallic salts from the water by means of soap takes time and is expensive, and other substances are com- monly used for this purpose. If the water is temporarily hard it can be softened, as we have seen, by boiling, but if large amounts of water are to be used in a laundry the cost of heating the water is prohibitive. Water containing bicarbonates can be softened by adding to it just enough calcium hydroxide to convert the acid salts into neutral salts: Ca(HCO 3 ) 2 + Ca(OH) 2 = 2CaCO 3 + 2H 2 An excess must be avoided, because calcium hydroxide is soluble and its presence in water renders it hard. Lime is employed to soften water to be used in boilers. Ammonium hydroxide is commonly used in the household; in this case the reaction which occurs is as follows: Ca(HCO 3 ) 2 + 2NH 4 OH = CaC0 3 + (NH 4 ) 2 CO 3 + 2H 2 O Permanently hard water may contain in addition to calcium sulphate the soluble salts of other metals; sea- water is hard because there is present in it, along with other substances, the chlorides of calcium and magnesium. In order to soften hard water of this type, it is treated with the sodium salt of an acid which forms insoluble calcium and magnesium salts. A carbonate or borate (borax) can be used for this purpose: CaSO 4 + Na 2 CO 3 = CaCO 3 + Na 2 SO 4 630. In the " permutit " process for softening water, the water is caused to flow slowly over an artificial silicate which approaches in composition the formula NaAlSi0 4 ,3H 2 O. Silicates of similar composition, which are derivatives of orthosilicic acid, H 4 SiO 4 , occur in nature and are called zeolites. They are much more readily attacked by water than feldspar, KAlSi 3 Og, and similar minerals which are richer in silica. Permutit is furnished in the form of flat granules about one-quarter of an inch in diameter. When it comes into contact with hard water a double decomposi- tion takes place, sodium passes into solution, and is replaced in the silicate by calcium and magnesium. When the permutit no longer acts in this way, it is treated with a strong solution of salt; the reaction is reversed and the material regenerated. The CALCIUM, STRONTIUM, BARIUM, AND RADIUM 537 process is a striking practical example of the application of the law of molecular concentration to reversible reactions. 631. The results of the analysis of a hard water are usually expressed in parts per million, the substances producing hardness being calculated as calcium carbonate. It is also expressed in " degrees," which in the United States indicate the number of grains per U. S. gallon: 1 degree equals 17.1 parts per million. Hard water from wells in limestone regions may contain as much as 400 parts per million. The action of hard water in boilers will be considered later. 632. Calcium Oxide. Quicklime, or lime, CaO, is a very important industrial product and is made in large quantities by heating limestone: CaC0 3 ^ CaO + C0 2 The reaction is a reversible one and the equilibrium attained is determined by the temperature, provided the gas does not escape. At 700 the pressure of carbon dioxide in equilibrium with the oxide is 25 mm. and at 900 it reaches 1 atmosphere. In making lime, calcium carbonate is " burned " at about 1000 in kilns 40 to 50 feet high and 6 to 10 feet in diameter, constructed of brick or of iron plates lined with fire-brick. The limestone is added at the top of the kiln and the lime withdrawn from the bottom. Since lime and calcium carbonate are poor conductors of heat the fuel selected for use is that which produces a flame which penetrates some distance into the kiln; wood, soft coal, oil, or gas are used for this purpose. An ingenious device is now used in certain kilns which markedly simplifies the lime-burning. The gases are drawn off from the top of the kiln by means of a fan and a part of them is forced into the air supply furnished the grate where the fuel is burning. Dilution of the air with carbon dioxide lowers the tem- perature of the flame (230) and increases its size; as the result of the " long-flame " combustion of the fuel a more even heating of the charge at a lower temperature is secured (233). The limestone used in making lime is more or less contam- inated with magnesium carbonate, silica, clay, oxides of iron, and other substances, and if it is heated during its formation above 1200, the calcium oxide and the silica react; fusible silicates are formed which fill up to some extent the pores in the calcium 538 INORGANIC CHEMISTRY FOR COLLEGES oxide. Such lime is said to be " over-burned " and reacts slowly with water. A lime which contains less than 5 per cent of impurities is called " fat " or " rich." Magnesium limes may contain as much as 30 per cent of magnesium oxide; they react with water slowly, and yield a mixture of calcium and magnesium hydroxides, which produce a " smooth " texture; limes of this kind are used in plasters for finishing purposes. 633. Slaked Lime. The addition of water to calcium oxide to form calcium hydroxide is called " slaking ": CaO + H 2 O = Ca(OH) 2 The union is accompanied by the evolution of heat, which in con- fined places may result in fire. In slaking lime to prepare mortar the former is treated with about twice the amount of water required to convert it into cal- cium hydroxide. The thick paste formed when the lime has slaked is mixed with 2.5 to 3 volumes of sand, and water is added to give the mortar the right consistency. When mortar sets it first loses the water it contains and then slowly absorbs carbon dioxide from the air; the calcium carbonate formed crystallizes around the grains of sand and binds them together. Calcium carbonate is formed to a depth of about 0.1 inch during the first year, but as it increases in thickness the change takes place more slowly. About twenty-three years are required to convert the mortar in an ordinary brick wall into cal- cium carbonate. The presence of sand in mortar renders the mixture more or less porous when it dries, and the absorption of carbon dioxide is, therefore, facilitated. Lime that is exposed to the air absorbs carbon dioxide, and when it is converted in this way into calcium carbonate it will not react with water; it is said to be " air-slaked." 634. Calcium Hydroxide. The way in which calcium hydrox- ide is prepared from lime has just been described (633). The compound finds many important uses in industrial chemistry because it is the cheapest base. It dissolves slightly in water (1 part in 600 parts at 18); the solution, called limewater, is used in medicine when a very dilute solution of an alkali is needed. A CALCIUM, STRONTIUM, BARIUM, AND RADIUM 539 suspension of the finely divided solid in water is known as " milk of lime." Calcium hydroxide is used in making the caustic alkalies and bleaching powder; for purifying illuminating gas; in the manu- facture of many chemicals, such as acetic acid; in purifying sugar solutions; in bleaching cotton; in tanning to remove the hair from hides, etc. 635. Calcium Sulphate. A hydrated form of calcium sulphate, CaSO4,2H 2 O, occurs as the mineral gypsum. It is mined in large quantities and is used in making plaster and stucco. A pure granular form of the dihydrate is called alabaster, and as it takes a high polish it is used in making statuary. The solubility of gypsum in water is 0.2 gram in 100 grams. When gypsum is heated it loses water and is first converted into a hemihydrate 2(CaSO 4 ,2H 2 O) ^ (CaSO 4 ) 2 ,H 2 O -f- 3H 2 O three-quarters of the water being lost. At about 200 the rest of the water is given off and the anhydrous form of the salt is obtained. The hemihydrate when mixed with water unites with it readily, whereas the union of anhydrous calcium sulphate and water takes place very slowly. 636. Plaster of Paris, which is the hemihydrate, (CaSO4) 2 ,H 2 O, is manufactured in large quantities by calcining gypsum at about 145 in muffles or rotary cylinders. When the finely ground material is mixed with about twice the theoretical quantity of water required to convert it into gypsum, it forms a plastic mass which "sets" in a short time to a solid. Since the mass expands on setting, plaster of Paris can be cast in molds, and for this reason it is used in making casts of statuary, wall decorations, etc. These are frequently treated with a solution of paraffin or stearin in gasoline, which leaves on evaporation a thin coating of the wax that fills the pores of the plaster. This treatment prevents the solution of the plaster by water and the accumulation of dust. Solutions of this kind are also valuable for water-proofing articles made of textile fabrics, such as tents, etc. The setting of plaster of Paris results from the fact that when it is mixed with water the latter dissolves some of the hemihydrate which unites with water to form the dihydrate. Since the latter 540 INORGANIC CHEMISTRY FOR COLLEGES is less soluble than the former it crystallizes from the solution, and more of the hemihydrate dissolves; the process continues until the plaster has been changed into a mass of interlocking crystals from which the excess of water evaporates. If the plaster has been "hard-burned" or " dead-burned " it requires hours or days for the setting, which takes place in the case of ordi- nary plaster of Paris in a few minutes; the product obtained from dead-burned plaster is hard and strong and is preferred for floors and for hard-finish plasters. The fact that the presence of colloidal substances retards crystallization is utilized when it is advisable to have the setting of plaster of Paris take place slowly. This is accomplished by adding to the plaster, glue, vegetable gums, or fine saw-dust; such a mixture is called " stucco." The name stucco is also applied to the mortar finish put on the exterior walls of buildings; it is a cement-sand mortar containing some fiber to give it strength. Wall plaster contains in addition to the retarding agent about 2 pounds of hair per ton, which give additional strength to the plaster. The precipitated calcium sulphate which is a by-product hi a number of chemical industries, is used under the names " crown filler " and " pearl hardening," as a filler in paper making and in weighting cloth. 637. The Phosphates of Calcium. The tertiary salt of ortho- phosphoric acid, Cas(P04)2, occurs widely distributed as phos- phorite and, in combination with calcium fluoride or chloride, as apatite, CaF2(Cl2),3Ca3(PO4)2. The ashes derived from bones contain about 80 per cent of calcium phosphate. The salt is precipitated when a soluble phosphate is added to a solution of a calcium salt; it is soluble in nitric acid (593), and is reprecipitated when the solution is neutralized. It is converted by sulphuric acid into calcium sulphate and the secondary salt, CaHPO4, the primary salt, Ca(H 2 PO4)2, or phosphoric acid, H 3 PO4, the extent of the conversion being determined by the relative amounts of the two substances used. The secondary salt is difficultly soluble in water (0.23 gram per liter); the primary salt is more soluble (18 grams per liter). Primary calcium phosphate is made in the pure condition for use in baking powders, and the mixture of the salt with calcium sulphate and some secondary phosphate is manu- CALCIUM, STRONTIUM, BARIUM, AND RADIUM 541 factured on the large scale for use as a fertilizer and sold under the name " super-phosphate." 638. Fertilizers. The importance of ammonia or nitrates and of potassium salts in the growth of plants has been emphasized. In addition to these substances a growing plant takes from the soil phosphates, silicon, iron, sodium, magnesium, chlorine, and other elements. Most soils contain an adequate supply of all these except phosphates, which are present in certain regions in limited amounts only. When any of the plant-foods become deficient they must be added to the soil if it is to retain its fertility. Nitro- gen is supplied as sodium nitrate, ammonium sulphate, calcium nitrate, calcium cyanamide (342), or in the form of guano, manure or sewage; the source of potassium is the chloride or sulphate, and that of phosphoric acid one or more of the calcium phosphates. A " complete "fertilizer contains substances which furnish nitro- gen, potassium, and phosphorus, but as many soils lack only one of the plant foods, the materials are used separately when the nature of the soil is known. The soil in regions that contain granitic rocks are usually rich in potassium, which is derived from feldspar; other soils are supplied with phosphates from the phos- phorite they contain. Lime or finely powdered calcium carbonate is added to soils which have become too acid as the result of the disintegration of the organic matter present, which forms the so-called " humus." Calcium sulphate is also added to acid soils. It is believed that the salt serves to fix the ammonia contained in the air. Ammonia and carbon dioxide form ammonium carbonate, which inter- acts with the sulphate to produce calcium carbonate and ammo- nium sulphate. The former neutralizes the organic acids, and the latter, being more stable than ammonium carbonate, remains in the soil and thus furnishes a supply of fixed nitrogen. Garbage is now treated in large cities to recover the valuable products it contains. It is heated with steam under pressure and the soft pulp produced is pressed ; the oil is separated and used for soap and candle stock, and the press cake is dried and used as a " filler " in fertilizers, since it contains nitrogen, phosphorus, and a little potassium. Bones which have been heated with steam to remove organic material are ground and used as fertilizers, but as the solubility 542 INORGANIC CHEMISTRY FOR COLLEGES of tricalcium phosphate is very small, bone-meal is not a good fertilizer. The organic matter contained in bones interferes with its solution in the soil, and it is at times removed by heating the bones to a high temperature in the presence of air. To render the material more soluble, it is usually converted into super- phosphate with sulphuric acid. The most important sources of phosphoric acid for fertilizers are the great beds of phosphorite which are more or less widely distributed. Those occurring in Florida, South Carolina, and Ten- nessee are placed so they are easy to mine, and because they are porous they are readily handled. The ground rock is treated with enough " chamber " acid to react according to the following equation: Ca 3 (PO 4 ) 2 + 2H 2 S0 4 + 6H 2 O = Ca(H 2 PO 4 )2,2H 2 O + 2(CaSO 4 ,2H 2 O) The mixture produced contains some phosphoric acid in addition to the primary salt; it is sold under the name of superphosphate. If the conversion of the tertiary salt has not been completely effected, the part left in the mixture reacts on standing with some of the primary salt to form the secondary salt: Ca 3 (PO 4 ) 2 + Ca(H 2 PO 4 ) 2 = 4CaHPO 4 A similar reaction takes place between the tertiary salt and the free phosphoric acid : Ca 3 (PO 4 ) 2 + H 3 PO 4 = 3CaHPO 4 These reactions constitute what is called " reversion." Since secondary calcium phosphate is very difficultly soluble in water, a " reverted " phosphate is absorbed much more slowly by the soil and is much less valuable as a fertilizer. 639. Bleaching Powder. This compound, which is also called chloride of lime or chlorinated lime, is manufactured by treating slaked lime with chlorine, either in absorption chambers built of brick, cast-iron, or lead, or in cast-iron cylinders provided with rotating blades which act as conveyors and keep the solid in motion so that it comes in contact with the gas. The chlorine must be admitted slowly as the reaction which takes place is exothermic and a rise in temperature interferes with the absorption. The CALCIUM, STRONTIUM, BARIUM, AND RADIUM 543 materials are left in contact until the bleaching powder formed contains from 36 to 37 per cent " available " chlorine, that is, chlorine which is liberated when the powder is treated with an acid. Bleaching powder is a mixed salt of hydrochloric and hypochlorous ,c\ acids and has the formula Ca4, prepared by adding a soluble chromate to a barium salt, is insoluble in acetic acid, whereas the chromates of calcium and strontium, which are insoluble in water, dissolve in the acid. Barium is determined quantitatively by precipitating and weighing it as barium sulphate. RADIUM 652. The remarkable properties of radium salts, the study of which has had such an influence on the development of our knowl- edge of the constitution of matter, will be discussed in some detail later; at this point attention will be drawn only to the properties of the metal and its compounds which resemble those of the other members of the calcium family. In 1896 Becquerel, a French scientist, discovered that minerals of uranium gave off a radiation that affected a photographic plate and passed through bodies opaque to light. In 1903, M. and Mme. Curie showed that the radiation produced by minerals containing uranium was stronger than that observed in the case of pure salts of the metal. It was concluded, therefore, that the minerals contained another substance more radio-active than uranium, and, as a result of a careful study of the problem, a new element, which was named radium, was discovered. The element was separated from the mineral pitch-blende, which contains uranium oxide, U2O3, along with a large number of compounds of other elements. When the constituents of the mineral were separated by the methods used in qualitative analysis, it was found that the barium sulphate obtained was the most active of all. This was converted into the oxide from which the bromide was prepared. Repeated fractional crystallization of this salt served to separate it into barium bromide and the bromide of the new element. In 1910 the metal was obtained in the free condition by elec- trolyzing a solution of radium chloride, using mercury as the CALCIUM, STRONTIUM, BARIUM, AND RADIUM 549 cathode. The solution of the metal prepared in this way was heated to remove the mercury by distillation. The radium obtained was found to be a white metal, which melted at 700, tarnished in the air, and decomposed water with the formation of the hydroxide and the evolution of hydrogen. Radium salts are now manufactured in the United States under the supervision of the Bureau of Mines from carnotite, a uranium ore which is found in Colorado. It is estimated that a ton of the minerals used as a source of radium contains about 0.2 gram of the element; it is for this reason that the cost of pure radium salts is so high about $120,000 per gram. Highly impure salts of radium can be used for most purposes to which it is put, and the cost of the material is correspondingly less. Many rocks and the waters of some mineral springs show the property of radio-activity to a slight degree. The spectrum of radium consists of two red bands, a blue-green band, and two faint violet lines. The properties of the salts of the metal, as far as they have been studied, are in accord with the position of the element in the periodic classification. EXERCISES 1. Summarize in the form of a table the physical and chemical properties of the alkaline earth metals and their oxides, chlorides, nitrates, carbonates, and sulphates. State what you think would be the properties of the analo- gous compounds of radium, taking into consideration the properties of the com- pounds of the other metals of the group and the position of radium in the family. 2. (a) Why is a very soluble salt more efficient in making freezing mix- tures than a less soluble one? (6) Should CaCl 2 , 6H 2 O or CaCl 2 be used in making a freezing mixture? Give a reason for your answer . 3. How could you show that dolomite is a carbonate which contains magnesium and calcium? 4. How could you distinguish powdered chalk, marble, and precipitated calcium carbonate from one another? 5. Write equations for the reactions which take place in the case of the fol- lowing: (a) Ca 3 N 2 +H 2 O, (6) NaHCO 3 + CaCl 2 , (c) CaCO 3 + C1 2 -f H 2 O, (d) NaOH -j- temporarily hard water, (e) NaOH + permanently hard water, and (/) NaOH + both temporarily and permanently hard water, (g) MgSO 4 + soap. 6. Write equations for the reactions which take place (a) in the soften- ing of water by permutit, and (6) the regeneration of the latter. 7. (a) What weight of sodium carbonate is required to soften 1,000,000 550 INORGANIC CHEMISTRY FOR COLLEGES gallons of permanently hard water which shows 20 degrees of hardness? (6) If a water is temporarily hard (20 degrees) what weight of lime is re- quired to soften 1,000,000 gallons of it? 8. (a) Write equations for the reactions by which tricalcium phosphate can be converted into the secondary and primary salts, and free phosphoric acid. (6) How could you readily distinguish from one another the three salts? 9. (a) Calculate the theoretical percentage of available chlorine in a pure sample of bleaching powder. (6) What percentage of this is obtained from commercial bleaching powder? Devise a method to determine (c) the total chlorine and (d) the available chlorine in a sample of bleaching powder. 10. Calculate the molar solubility of Ca(OH) 2 , Sr(OH) 2 , and Ba(OH) 2 in water at 18 and state the normality of a saturated solution of each. 11. Starting with BaSO 4 write equations for the reactions used in making the following: (a) BaCl?, (6) Ba(NO 3 ) 2 , (c) Ba 3 (PO 4 ) 2 , (d) BaO 2 . CHAPTER XXXVII BERYLLIUM, MAGNESIUM, ZINC, CADMIUM, AND MERCURY 653. The elements in the second family of the second group in the periodic classification do not show to a marked degree the gradation in physical and chemical properties which is ordinarily observed in the case of other families. In their compounds they are bivalent, but mercury forms also salts in which the metal has the valence 1. The hydroxides of the metals are bases which are less active than those of the calcium family, and the salts derived from them are more or less hydrolyzed by water. The hydroxide of beryllium, which is also called glucinum, and that of zinc are weakly acidic and dissolve in solutions of the caustic alkalies. Beryllium is a rare element; it occurs as a double metasilicate with aluminium in the mineral beryl, Al2Be3(SiOs)6, which is called emerald when colored green by the presence of a small amount of chromium silicate. Beryllium carbonate dissolves in water and is highly hydrolyzed in solution. MAGNESIUM 654. The compounds of magnesium are widely and abun- dantly distributed in nature. Dolomite, a double carbonate of magnesium and calcium, occurs as a rock and is the chief constit- uent of certain mountain ranges. Magnesite is magnesium car- bonate, talc or soapstone (510) is a hydrated silicate of mag- nesium, and asbestos (510) is an anhydrous silicate. The sulphate and chloride of magnesium occur in sea water and in salt beds. Compounds of magnesium have been used for many years. The sulphate occurs in the water of the mineral springs at Epsom, England, and for centuries has been used under the name Epsom salt. The basic carbonate, " magnesia alba," was known to the 551 552 INORGANIC CHEMISTRY FOR COLLEGES alchemists. The free element was first isolated in an impure condition by Davy, and was prepared in 1830 by Liebig, by the action of potassium on fused magnesium chloride. Magnesium is manufactured by the electrolysis of a fused mixture of magnesium and potassium chlorides, which are obtained from carnallite, MgCl2,KCl,6H2O. The preparation is carried out in an iron vessel which serves as a cathode, the anode being a rod of carbon. Magnesium is a silver-white metal, which is characterized by having the very low specific gravity 1.75; it melts at 650, boils at 1120, and can be drawn into the form of a wire and rolled into sheets. The metal on account of its lightness has recently been employed in making alloys to be used in the construction of aero- planes (528). Magnesium is less active than the metals of the calcium family; it tarnishes in air and reacts slowly with water. The magnesium-mercury couple (545) made by treating the metal with a little mercuric chloride decomposes water fairly rapidly. Magnesium burns in the air with an intense white light which is rich in the shorter wave lengths of the visible spectrum (blue and violet) ; it also emits the so-called ultra-violet rays, which are still shorter in wave lengths, are not visible, but affect the materials used in photographic plates. The light of burning magnesium is said to be about sixty times as active chemically as that produced by burning carbon. In photography magnesium is burned directly in the form of a thin sheet or as a powder. In the latter case the finely divided metal is mixed with powdered potassium chlorate, the amount of the salt used being that required to furnish just enough oxygen to convert the metal into oxide. When such a mixture is ignited it burns instantaneously, whereas the powder burns sluggishly and soon becomes coated with oxide, which stops the reaction. When magnesium burns in the air a part of the metal is converted into the nitride, MgsN2, which is decom- posed by boiling water with the formation of magnesium hydroxide and ammonia. 655. The Oxide and Hydroxide of Magnesium. Magnesium oxide is prepared by heating magnesite, MgCOs, at about 1700; it is known as calcined magnesia and is obtained as a white porous powder, which is infusible and does not react with water. It is employed in the manufacture of fire-brick that are used in lining BERYLLIUM, MAGNESIUM, ZINC, CADMIUM, AND MERCURY 553 furnaces in which reactions are carried out at very high tempera- tures. The oxide prepared by heating the precipitated carbonate is obtained as a light powder; it is used in the manufacture of coverings for steam-pipes as an insulator to prevent the transfer of heat. It reacts slowly with water to form magnesium hy- droxide. Magnesium hydroxide, Mg(OH) 2 , is made by treating a mag- nesium salt with a solution of an alkali. It is very slightly soluble in water, about 0.01 gram dissolving in 1 liter, but the concentra- tion of a saturated solution is great enough to permit the reaction of the base with ammonium salts, which, accordingly, dissolve it: Mg(OH) 2 + 2NH 4 C1 = MgCl 2 + 2NH 3 + 2H 2 The fact that the hydroxide is not precipitated by ammonium hydroxide in the presence of ammonium salts has already been explained at some length (589) . 656. Magnesium Chloride. This salt resembles calcium chlo- ride in composition and behavior. It crystallizes from water as a hexahydrate, MgCl 2 ,6H 2 O, and is very soluble and very hygro- scopic. Like the chloride of calcium it cannot be dehydrated without decomposition, hydrolysis taking place even to a greater extent than in the case of the calcium salt. Anhydrous mag- nesium chloride can be prepared by first dehydrating a double salt of the composition NH4Cl,MgCl 2 ,6H 2 O, and then driving off the ammonium chloride at a high temperature. The hydrolysis of magnesium chloride in boiling water is appreciable, MgCl 2 + 2H 2 ^ Mg(OH) 2 + 2HC1 and for this reason sea-water and other water which contains this salt cannot be used in boilers. 657. Magnesium Sulphate. A number of hydrates of this salt are known; the one having the composition MgSO4,7H2O is obtained when a solution of magnesium sulphate is evaporated so that crystals are formed from it at room-temperature. It is efflorescent and very soluble in water. It is used in dyeing, cotton printing, and other industries, and in medicine as a purgative. The monohydrate of magnesium sulphate, MgSO4,H 2 O, ap- proaches in solubility that of gypsum, CaS04,2H 2 O, being almost 554 INORGANIC CHEMISTRY FOR COLLEGES insoluble in water; it occurs in the salt beds at Stassfurt as the mineral kieserite. 658. Magnesium Carbonate. This compound occurs as the mineral magnesite, MgCOs, and as a constituent of dolomite (626). It is not formed in the pure condition by precipitation when a soluble carbonate is added to a magnesium salt, the product obtained being either a mixture of magnesium carbonate and magnesium hydroxide or a basic salt of the metal. The product obtained by the precipitation is called " magnesia alba," and is used in medicine and as a cosmetic. 659. Boiler Scale. The composition of hard water and the methods employed in its treatment when it is to be used for industrial purposes have been discussed (629). It was stated that it should not be used in boilers for the production of steam; this is due to the fact that it produces on the boiler tubes a deposit which is a poor conductor of heat, and the fact that acids are formed which corrode the metal of which the boiler is constructed. The solids which separate in a boiler when hard water is evap- orated to make steam, are produced either as a porous, non-adher- ing material, which can be easily removed by " blowing off " or washing out the boiler, or as a dense solid which coats the tubes and sides and forms what is called " boiler scale." In the first case the material is produced largely as the result of the con- version of the soluble bicarbonates present into the carbonates, which precipitate. The scale is composed principally of calcium sulphate and the monohydrate of magnesium sulphate, MgSO^EbO (657) , which is practically insoluble and is formed from the soluble sulphate at the temperature obtained in the boiler. Boiler scale- also contains magnesium hydroxide which is present in it as the result of the hydrolysis of the magnesium chloride dissolved in the water. It will be recalled that calcium sulphate is less soluble in hot water than in cold water; the solubility drops off rapidly as the water is heated to the high temperatures obtained in a boiler. At 68 F. water dissolves 140.6 grains of the salt per gallon; at 212 F. the solubility is reduced to 125.9 grains, at 284 F. when the steam pressure is 37 pounds the solubility is 45.6 grains, and at 356 F. which corresponds to a steam pressure of 131 pounds it is 10.5 grains per gallon. The heat conductivity of the compounds present in boiler scale BERYLLIUM, MAGNESIUM, ZINC, CADMIUM, AND MERCURY 555 is small compared with that of iron; that of calcium sulphate has been found to be about one-fiftieth that of the metal. As a result, when a boiler the surface of which is covered with scale is used to make steam, it becomes overheated and the metal may get red hot and become deformed. As the tubes become filled with scale the surface exposed to the water grows smaller and the steam-producing efficiency of the boiler rapidly falls off. The corrosion of boilers by hard water is due chiefty to the presence of magnesium chloride, which hydrolyzes appreciably at the high temperatures reached. The acid formed dissolves the iron of the boiler. 660. Tests for Magnesium. The fact that sodium carbonate precipitates a colorless, insoluble, basic carbonate, which is not formed in the presence of ammonium salts, serves to separate magnesium from the metals of the alkaline earths. Magnesium- ammonium phosphate, MgNH4PO4,OH2O, is formed as a crys- talline precipitate when disodium phosphate, ammonium hydrox- ide, and ammonium chloride are added to a solution of a mag- nesium salt. The chloride is added to prevent the precipitation of magnesium hydroxide (589) by the ammonium hydroxide. Magnesium-ammonium phosphate is formed and used in the quantitative determination of magnesium or of phosphates. When it is ignited it is converted into magnesium pyrophosphate, Mg2P20r. Magnesium salts give no color to the Bunsen flame. ZINC 661. The chief ores of zinc are smithsonite or zinc spar, ZnCOs, and sphalerite or zinc blende, ZnS; zinc silicate, which occurs as calamine, Zn2Si04,H2O, and franklinite, Zn(FeO2)2, are also used. The metal is obtained, after roasting the ore, by reduction with finely divided coal in cylindrical fire-clay retorts, 4 to 5 feet long and 8 to 10 inches in diameter, which are heated by gas. Since reduction of the oxide takes place at a temperature above the boiling-point of the metal, the vapor of the latter which issues from the retort is led through a condenser made of fire-clay, where it is deposited in the molten state. At first a mixture of zinc oxide and metallic zinc in the form of a fine powder, which is known as zinc dust, is deposited ; it contains about 10 per cent zinc oxide and on account of the fine state of the metal finds important uses as a 556 INORGANIC CHEMISTRY FOR COLLEGES reducing agent. The liquid zinc drawn off from time to time is cast in molds and is called " spelter "; it usually contains lead, cadmium, and iron, and smaller quantities of arsenic, antimony, and sulphur. It is refined by melting in a furnace; the lead pres- ent combines with a little zinc and settles to the bottom, and on this is deposited a zinc-iron alloy. The purest zinc is made by distillation at as low a temperature as possible. Zinc melts at 419, boils at 940, and has the specific gravity 7.08 when cast. Cast zinc is brittle, but when heated at 120 it can be rolled out into sheets, which remain pliable when cooled. Just above 200 it becomes so brittle it can be powdered. The so-called " granulated " zinc used in the laboratory is prepared by dropping the molten metal into water. In moist air zinc becomes covered with a thin coating of basic carbonate, which protects it from further action (550). It burns in air when heated, reacts with most non-metals, forms alloys with other metals (542), decomposes steam at moderately high temperatures, and reacts with hot solutions of alkalies (554). The effect of impurities in the metal on its solution in acids has been described (567). Zinc is used for roofs, gutters, etc., in the manufacture of brass and other alloys, and as a protective coating for iron which is "galvanized" or " sherardized " (551). Large quantities are consumed in dry cells (568). The metal contracts but little in passing from the liquid to the solid state and can be cast; so-called French bronzes are cast zinc, which is electroplated with bronze or brass. 662. Zinc Oxide. This compound is used extensively under the name zinc white or Chinese white. It is manufactured by heating metallic zinc and conducting the vapor of the metal through a flue to which air is admitted; the oxide formed from the burning metal is collected after cooling by passing it along with the excess of air and nitrogen through settling chambers and large bags of cotton cloth. At times the vapor of the metal is obtained directly as the result of the reduction of the ore with coal. Zinc oxide turns yellow when hot, but regains its whiteness on cooling. It is used as a pigment in making paint and as a filler in the manufacture of automobile tires. On account of its anti- septic properties it is used in ointments. BERYLLIUM, MAGNESIUM, ZINC, CADMIUM, AND MERCURY 557 663. Zinc Hydroxide. The hydroxide of zinc, which is formed on the addition of a soluble hydroxide to a zinc salt, shows both basic and acidic properties. Its salts are hydrolyzed; zinc sul- phate shows an acidic reaction and sodium zincate, Na2ZnO2, a strong basic reaction. In order to convert zinc hydroxide into the latter salt an excess of sodium hydroxide must be used over that indicated by the following reaction; Zn(OH) 2 + 2NaOH ^ Zn(ONa) 2 + 2H 2 O Zinc hydroxide dissolves in a solution of ammonia as the result of the formation of the compound Zn(NH3)4(OH)2 (595). 664. Zinc Chloride. A number of hydrates of zinc chloride are known, all of which are deliquescent. The salt is usually sold in the anhydrous condition; it melts at 290 and boils at 730, and is hydrolyzed by water. When mixed with zinc oxide and a little water it forms a plastic mass which sets to a solid: ZnCl 2 + ZnO = Zn 2 OCl 2 The product is used as a cement and for other purposes. Zinc chloride is used as a flux in soldering. 665. Other Salts of Zinc. Zinc carbonate is formed as a pre- cipitate when sodium bicarbonate is added to a solution of a zinc salt; sodium carbonate which is slightly hydrolyzed produces a basic salt Zn(OH) 2 ,ZnCO3. Zinc sulphate, ZnSO4,7H 2 O, is some- times called white vitriol; it is isomorphous with green vitriol, FeSO4,7H 2 O, and the sulphates of other bivalent elements which crystallize with 7 molecules of water. Zinc sulphide is a con- stituent of lithophone (649). 666. Test for Zinc Salts. The sulphide of zinc, ZnS, is formed as a white precipitate when ammonium sulphide is added to a solution of a zinc salt or a zincate. It is not precipitated by hydrogen sulphide in the presence of strong acids, in which it is soluble. Zinc hydroxide is precipitated by sodium hydroxide and dissolves in an excess of the reagent; it dissolves also in ammonium hydroxide. In determining zinc quantitatively it is usually pre- cipitated as the carbonate, which is ignited and weighed as oxide. 558 INORGANIC CHEMISTRY FOR COLLEGES CADMIUM 667. Cadmium is obtained as a by-product in zinc smelting. The small amount of the metal present in the ores is found in the zinc dust which is first condensed from the vapors as they issue from the retorts (661). The metal is more volatile than zinc and is separated by distillation from the zinc dust after carbon has been added to reduce the oxides present. Cadmium is a silver-white metal which melts at 321 and boils at 778. It dissolves in acids less readily than zinc. It is used as a constituent of low-melting alloys (492) , and in amalgams for dental use. Cadmium sulphide, which is bright yellow, is employed as an artist's color; the bromide and iodide are used in precipitating the silver halides that are the active constituents of photographic plates and films. 668. Compounds of Cadmium. A large number of com- pounds of the metal are known; only a few need be mentioned here. The chloride, CdCl2,2H2O, is efflorescent and is not hydro- lyzed by water. The iodide, Cdl2, is but slightly ionized in solu- tion a fact thought to be due to the formation of a salt that is composed of two molecules of the iodide, to which the formula Cd(Cdl4) is given. This view is a reasonable one since cadmium is characterized by its ability to form stable complex salts. Among these are double chlorides and double cyanides. Cadmium sul- phate, 3CdSO4,8H20, is not isomorphous with the sulphates of zinc and magnesium. The precipitated carbonate of the metal is a normal salt. 669. Tests for Cadmium Salts. Cadmium hydroxide is formed as a white precipitate when a solution of an alkali is added to a cadmium salt; it is insoluble in alkalies, but dissolves in ammonium hydroxide to form the compound Cd(NHs)4(OH)2. The sul- phide, which is yellow, is insoluble in dilute solutions of strong acids, and is, consequently, precipitated along with copper and the other metals of the so-called second group in qualitative analysis. It dissolves in strong solutions of acids. Cadmium can be sep- arated from copper by adding potassium cyanide to the solution of the compounds of the two metals and passing in hydrogen sul- phide. The soluble complex cyanide of copper which is formed, KCu(CN)2, is not decomposed, whereas the cadmium salt, BERYLLIUM, MAGNESIUM, ZINC, CADMIUM, AND MERCURY 559 K2Cd(CN)4, is converted into cadmium sulphide, which appears as a yellow precipitate. MERCURY 670. On account of its occurrence in the free condition mer- cury has been known for over 2000 years. The fact that it is a liquid and resembles silver in its properties led to the Latinized name of the element, hydrargyrum, which was derived from the Greek words signifying liquid silver. Mercury played an impor- tant part in the early theories of the constitution of matter. It was named by the alchemists after Mercury, the messenger of Jove, on account of its rapid disappearance when heated, and was indi- cated in alchemical books by the symbol which represented a messenger's wand the same symbol as that used in astrology to represent the planet Mercury. The salts of mercury have a marked effect on the human body, and were studied in consider- able detail in the sixteenth century by the iatrochemists, who investigated the action of inorganic substances as drugs. Mercury forms two classes of compounds; the members of one of these, in which the metal is bivalent, resemble to some extent the analogous compounds of cadmium and zinc; those of the other class contain mercury with the valence 1. The compounds of univalent mercury possess properties analogous to those of the compounds of other heavy metals, copper and silver, when the latter have the valence 1 ; they are called mercurous compounds, and those derived from bivalent mercury are mercuric com- pounds. 671. Preparation and Properties of Mercury. The element is found in the free condition but is obtained chiefly from cinnabar, HgS, which occurs as a red crystalline mineral in California, Spain, Italy, and Austria. The ground mineral is heated in a stream of hot air which oxidizes it HgS + 62 = Hg + S(>2 and the vapor of the metal formed is condensed. It can be purified by distillation, or by bringing it into contact with dilute nitric acid, which dissolves the impurities present. The latter process is carried out in the laboratory by allowing the metal to fall in minute globules through the acid contained in a long glass tube. It can also be purified by bubbling air through it for a long time. 560 INORGANIC CHEMISTRY FOR COLLEGES The metals dissolved in the mercury are oxidized and form on its surface a scum, which can be removed by filtering the liquid through chamois skin. Mercury is the only metal which is liquid at ordinary tempera- tures; it melts at -38.9, boils at 357, and has the specific grav- ity 13.6 at 0. It is used in barometers, thermometers, and air pumps for laboratory use, as a solvent for gold and silver in extract- ing these metals from their ores, as the cathode material in the electrolytic preparation of the caustic alkalies, in amalgamating zinc to be used in electric batteries (567), and in the preparation of amalgams with tin, silver or gold which are used by dentists to fill cavities in teeth. Finely divided mercury which has the appearance of a powder is used in medicine in the form of pills. Mercury unites with oxygen when it is heated with it at about 300 and forms mercuric oxide. It combines with the halogens and sulphur, does not react with dilute solutions of non-oxidizing acids, but dissolves in nitric acid. It decomposes hydrogen sulphide and is converted into black mercuric sulphide. Mercury dissolves many metals, with some of which it unites chemically. The compound of mercury and sodium having the composition NaHg2 is a solid. 672. The Oxides of Mercury. Mercurous oxide, Hg2O, is formed as a black precipitate when sodium hydroxide is added to a solution of a mercurous salt; it is slowly decomposed into mercury and oxygen by light or by gentle heating. Mercuric oxide, HgO, is formed in a similar way from mercuric salts; under these con- ditions it appears as a yellow, amorphous powder which turns dark brown when heated. When the oxide is prepared by heating mercury at 300-350, or by calcining mercuric nitrate, Hg(NO3)2, it is a bright red powder. No hydroxides of mercury exist; the reactions which would lead to their formation produce the oxides of the metal. 673. The Nitrates of Mercury. If mercury is dissolved in hot concentrated nitric acid it is converted into mercuric nitrate, Hg(NO3)2,8H2O; when left in contact with the dilute acid it reacts slowly to form mercurous nitrate, which crystallizes from water as a monohydrate, HgNO3,H2O. Both salts are hydrolyzed by water and form insoluble basic salts; they remain in solution only in the presence of an excess of nitric acid. Mercuric nitrate can be con- BERYLLIUM, MAGNESIUM, ZINC, CADMIUM, AND MERCURY 561 verted into mercurous nitrate by shaking its solution with mer- cury: Hg(N0 3 ) 2 + Hg = 2HgN0 3 674. The Halides of Mercury. Mercuric chloride, HgCl 2 , which is known as corrosive sublimate, is made by subliming a mixture of mercuric sulphate and sodium chloride: HgSO 4 + 2NaCl = HgCl 2 + Na 2 S0 4 It boils at 307, and when the mixture is heated to this temperature the conditions are such that the double decomposition takes place. Mercuric chloride unites directly with protein material and is, as a consequence, a poison. When taken into the body it interferes with the proper functioning of the kidneys, and the waste products formed as the result of the normal changes which take place, are not excreted and may finally produce death. The antidote com- monly used for corrosive sublimate is white of egg, which is a pro- tein. If this is administered soon enough it unites with the salt to form an insoluble compound, which can be removed from the stomach by means of an emetic. A very dilute solution of mer- curic chloride is used as an antiseptic in surgery. Mercurous chloride, HgCl, is used in medicine under the name calomel. It is prepared by subliming a mixture of mercuric chlo- ride and mercury HgCl 2 + Hg = 2HgCl or a mixture of mercuric sulphate, mercury, and salt. It is used in medicine as a purgative. Mercurous chloride is formed as a white precipitate w r hen a soluble chloride is added to a solution of a mercurous salt. It is also readily formed by reducing mercuric chloride; when stannous chloride, SnCl 2 , is used, it is oxidized to stannic chloride, and mercurous chloride or mercury is formed depending upon the quantity of the reducing agent used : 2HgCl 2 + SnCl 2 = 2HgCl + SnCU 2HgCl + SnCl 2 = 2Hg + SnCU The mercury appears as a black precipitate. 675. Mercuric iodide, HgI 2 , is obtained as a scarlet precipitate when a soluble iodide is added to a solution of a mercury salt. When the salt is heated above 128 it is converted into a form which is yellow, melts at 223, and on cooling yields yellow crystals. 562 INORGANIC CHEMISTRY FOR COLLEGES When these are cooled to room temperature they are unstable and change on being touched with a sharp object into the scarlet variety. Mercuric iodide dissolves in an excess of potassium iodide to form a stable, colorless, complex salt of the composition K 2 HgI 4 , from which the mercury is not precipitated by bases. Mercurous iodide, Hgl, is formed as a dark-green solid when iodine is rubbed with an excess of mercury; when a soluble iodide is added to a solution of a mercurous salt it is precipitated as a greenish-yellow powder, which decomposes spontaneously into mercuric iodide and mercury: 2HgI = Hg + Hgl2. 676. When a solution of ammonia is added to mercury salts the reactions which take place are of a different nature from any that have been described; the so-called ammono-compounds are formed, of which those derived from the halides are of the most importance because they are used in analytical chemistry. The addition of a solution of ammonia to one of mercuric chloride causes the precipitation of a white compound of the formula Hg(NH2)Cl which is called, "infusible white precipitate" or, better, ammono-basic mercuric chloride: HgCl 2 + 2NH 3 = Hg(NH 2 )Cl + NH 4 C1 The interaction of ammonia to form a compound of this type as the result of the replacement of a part of the acid radical by the NH 2 group resembles closely reactions of hydrolysis; in the latter case water separates into H and OH, the hydrogen unites with the acid radical, and the OH takes the place of the latter and a basic salt is formed : HgCl 2 + HOH = Hg(OH)Cl + HC1 in the case of ammonia the separation of the latter into H and NH 2 takes place and the hydrogen, as before, unites with a part of the acid radical, which is replaced by the NH 2 group. On account of this close analogy the reaction is said to be one of ammonolysis and the salt formed is called an ammono-basic salt. 677. When mercurous chloride is treated with a solution of ammonia, it decomposes into mercury and mercuric chloride and the latter reacts with the ammonia to form ammono-basic mercuric chloride. The mercury is deposited throughout the precipitate as a fine powder and, as a result, the precipitate is black. It was BERYLLIUM, MAGNESIUM, ZINC, CADMIUM, AND MERCURY 563 this reaction that gave its name to calomel, the word being derived from the Greek words which signify " beautiful black." 678. In the case of mercuric iodide the reaction is more com- plex; it is probable that the compound Hg(NH2)I is first formed and then reacts with a second molecule of mercuric iodide as follows: Hg(NH 2 )I + HgI 2 - Hg(NHg)I + 2HI The two hydrogen atoms of the NH 2 group unite with the two iodine atoms in mercuric iodide to form hydrogen iodide, and are replaced by a mercury atom; the graphic formula of the compound is, accordingly, Hg = N Hg I. The reaction is brought about by adding ammonia to a solution of the iodide in potassium iodide (K2Hgl4). The substance is a brown precipitate and has high tinctorial power, that is, an exceedingly small amount of it can be recognized by its color. A solution of mercuric iodide in potassium iodide to which potassium hydroxide has been added is known as Nessler's reagent and is used to detect small amounts of ammonia in water analysis. If the water contains ammonium salts the base present in the solution liberates ammonia, which then interacts with the mer- curic iodide to produce a brown color. The amount of ammonia present in the solution is determined by matching the shade pro- duced by it with that formed in one of a series of standards pre- pared from the reagent and solutions containing known amounts of ammonia. This is an example of a so-called " colorimetric " method of analysis, many of which are used in quantitative analysis. 679. Sulphides of Mercury. When hydrogen sulphide is passed into a solution of a mercurous salt at -ordinary tempera- tures, a mixture of mercury and mercuric sulphide, HgS, is obtained. Mercurous sulphide, Hg 2 S, is stable only at low tem- peratures. Mercuric sulphide is formed as a black precipitate when hydrogen sulphide is passed into a mercuric salt. It is very insoluble in water and acids and is not attacked by boiling nitric acid, its inertness in this respect being utilized in separating it from the sulphides of other metals precipitated in the course of a qualitative analysis. It is soluble in aqua regia. The precipitated, black sulphide is converted into the bright red crystalline variety when it is sublimed. The sulphide, which 564 INORGANIC CHEMISTRY FOR COLLEGES occurs as cinnabar, was formerly much used as a red pigment under the name vermilion. The material used at present is manufac- tured either by heating a mixture of mercury and sulphur until the black sulphide is formed, and then subliming the product, or by heating a mixture of mercury and sulphur with a strong solu- tion of caustic potash. 680. Other Salts of Mercury. Mercuric cyanide, Hg(CN)2, is prepared from precipitated mercuric oxide and hydrocyanic acid, HCN; it forms prismatic crystals that decompose on heating into mercury and cyanogen, (CN)2, a poisonous gas, which burns with a characteristic flame composed of a pink cone surrounded by a blue mantle. Mercuric fulminate, Hg(ONC)2, is formed when mercury is dissolved in nitric acid and alcohol is added to the solution. It explodes rapidly when struck and is used in detonating caps to explode gunpowder, nitroglycerine, etc. 681. Tests for Mercury Salts. Mercurous salts are precip- itated by hydrochloric acid, and when the white mercurous chloride formed is treated with ammonia it turns black (676). Mercuric salts are precipitated by hydrogen sulphide in the presence of hydrochloric acid. The black sulphide is separated from other sulphides by treatment with hot nitric acid, in which it alone does not dissolve. It is then dissolved in aqua regia, and the solution of the chloride is treated with stannous chloride, which precipitates white mercurous chloride and finally mercury as a black powder (674). All the salts of mercury are volatile when heated. EXERCISES 1. How could you prepare MgSO 4 ,H 2 O from MgSO 4 , 7H 2 O? 2. Explain why Na 2 CO 3 precipitates CaCO 3 from calcium salts and a basic carbonate from magnesium salts. 3. A sample of a mineral containing only MgCOs and CaCOs, weighing 1 gram, was analyzed for Mg. The Mg2P2O 7 obtained weighed 0.6032 gram. Calculate (a) the weight of MgCO 3 in the sample, (6) the percentage of Mg in the mineral, (c) the weight of CO 2 that could be obtained from 1 gram of the mineral. 4. Would the addition of calcium hydroxide to a water which caused the corrosion of a boiler be advantageous? Why? 5. Show by ionic equations why the neutral carbonate of zinc is pre- cipitated by NaHCOs and the basic carbonate by BERYLLIUM, MAGNESIUM, ZINC, CADMIUM, AND MERCURY 565 6. (a) Write ionic equations to represent the hydrolysis of MgS and CaS. (b) In which case is the hydrolysis greater? Give a reason for your answer. 7. Write equations for reactions by means of which (a) magnesium and zinc salts can be distinguished in three different ways, and (6) zinc salts from cadmium salts in two ways. 8. What chemical reactions would occur if a sample of mercury contain- ing small amounts of zinc and lead were agitated with a solution of mercurous nitrate? For what purpose could the process be used? 9. Write equations for the reactions which take place between Hg and (a) concentrated HNO 3 and (&) dilute HNO 3 . 10. Write equations for reactions by which Hg(NO 3 ) 2 can be converted into (a) HgCl, (&) HgO, (c) HgCl 2 , (d) HgNO 3 , (e) HgSO 4 . CHAPTER XXXVIII ALUMINIUM 682. Aluminium, which is sometimes called aluminum, is the typical metal in the third group in the periodic classification of the elements. With the exception of boron, which shows acid- forming properties only, the other elements in the group resemble aluminium closely. Thallium, in addition to being trivalent, forms a number of compounds in which it shows the valence 1; these resemble closely the analogous derivatives of silver. Other metals which can show a higher valence than 3 and, therefore, fall into other groups in the periodic classification, form important compounds in which the element is trivalent; among these are chromium, iron, and manganese. The compounds derived from these metals in the trivalent condition resemble closely in their chemical properties the analogous compounds of aluminium. We shall learn, therefore, through the study of aluminium compounds the chemical behavior of the derivatives of a trivalent metal. A marked difference exists between the chemical properties of the compounds of the univalent metals and those of the analogous compounds of the bivalent metals. We have seen, for example, that the hydroxides of the alkali metals are strong bases, which are very soluble in water, and form salts that are not hydrolyzed. The carbonates of these metals resist high temperatures, and the salts of other oxygen acids are relatively stable when heated. The hydroxides of the bivalent metals are weaker bases, much less soluble in water, and their salts are hydrolyzed to a small extent. The carbonates of these metals are converted by heat into oxides, and the salts of other oxygen acids are much more readily decom- posed than those of the alkali metals. The decreased activity of the metals as base-forming elements when they are in the trivalent condition is marked. The hydrox- 566 ALUMINIUM 567 ides of these metals are very weakly base-forming, and the salts derived from them are highly hydrolyzed; carbonates cannot be prepared in aqueous solution by precipitation, because they are completely hydrolyzed by water. The salts are decomposed by heat at comparatively low temperatures. The hydroxides of some of the metals dissolve in solutions of strong bases, and act, therefore, as acids. A comparison of the behavior of the deriva- tives of the three types of metals brings out clearly the fact already mentioned, that, in general, low valence toward negative atoms or groups is associated with base-forming properties and high valence with acid-forming properties. 683. The occurrence of aluminium has already been noted. It is the most abundant of the metals and the most widely distributed. It is found in feldspars, micas, kaolin, clay, bauxite, cryolite, alun- ite, corundum, and certain gems (510). Compounds of aluminium have been known for many years, and they were recognized as being derived from a metal that had not been isolated. Many unsuccessful attempts were made to separate it from its oxide, but, finally, Wohler in 1827 obtained the metal as the result of the action of potassium on fused alu- minium chloride. This method was too expensive to be used for commercial purposes, and aluminium, which has found a very important place in industry, was not manufactured until the electrolytic method devised in 1886 by Hall, was put into success- ful operation. 684. Preparation and Uses of Aluminium. In the Hall process the metal is obtained by the electrolysis of aluminium oxide dis- solved in a bath of fused cryolite, AlFs,3NaF. The material is contained in rectangular iron pots that are lined with thick plates of carbon, which serve as the cathode. The anodes are large graphite rods. In starting the reduction, the anodes are brought into contact with the bottom of the tank and heat is developed by the current as the result of the poor contact established. Cry- olite is placed in the cell, and when the bath is in the molten con- dition the anodes are withdrawn from the cathode, and aluminium oxide is added. The resistance of the cell produces enough heat to keep the material in the fused condition. From time to time molten aluminium is withdrawn from the bottom of the cell, and more oxide is added; the process is a continuous one because 568 INORGANIC CHEMISTRY FOR COLLEGES the cryolite is not decomposed by the current and serves only as a solvent for the oxide. Aluminium has a very low density, 2.6; it is used in construc- tion when a metal is required and weight is an important consider- ation. It is ductile, malleable, and can be rolled. Its tensile strength is low in comparison with that of iron; it cannot be machined and polished readily, and does not yield good castings. These defects can be overcome by alloying it with other metals. Alloys of copper and aluminium which contain from 5 to 10 per cent of the latter are called aluminium bronzes. They have a fine yellow color resembling gold and are used in making imitation jewelry and statuary. The alloys which con- tain from 11 to 89 per cent aluminium are brittle and highly crystalline and are not used. Those that contain 90-93 per cent aluminium are used in making castings and are silver-white. On account of its low electrical resistance aluminium is used in certain cases in wires and cables as conductors. In a finely divided condition it is used in making a paint, and, in the form of leaf, for stamping letters on book covers. At high temperatures aluminium is an active deoxidizing agent. When a mixture of the powdered metal and a finely divided oxide of iron is ignited, a vigorous reaction takes place and aluminium oxide and iron in the molten condition are formed. The tem- perature reached is about 2300. The mixture, which is called thermite, is used in welding together the ends of steel rails, for mending broken shafts of marine engines, etc. It was used in incendiary bombs during the recent war. Oxides of metals which do not react with carbon at the temper- ature obtained by burning coal, are reduced when heated with alu- minium. The method has been used to prepare chromium and similar metals. Large quantities of aluminium are used in deoxidizing iron and steel; the metal unites with the gases in the metal, and castings free from blow holes are obtained. The chief physical and chemical properties of the metal have already been described. It dissolves in hydrochloric acid but not in dilute sulphuric or nitric acid. It is soluble in solutions of the caustic alkalies as the result of the formation of aluniin- ates (654). ALUMINIUM 569 685. Aluminium Hydroxide. When a salt of aluminium is treated with a solution of a base, aluminium hydroxide is formed as a gelatinous precipitate: A1CU + SNttiOH = A1(OH) 3 + 3NEUC1 The hydroxide dissolves in solutions of caustic alkalies as the result of the fact that it shows weakly acidic properties. The salts are derived from a dehydrated form of the hydroxide A1(OH) 3 - H 2 O = HAlO-j and are called metaluminates: NaOH + A1(OH) 3 = NaAlO 2 + 2H2O The salt is highly hydroly zed and remains in solution only in the presence of an excess of sodium hydroxide. When an acid is added cautiously to the solution, aluminium hydroxide is repre- cipitated. The hydroxide does not dissolve in ammonium hydrox- ide, which is, accordingly, used when it is desired to precipitate it quantitatively. 686. Aluminium Oxide. Corundum, AfeOs, is a pure crys- talline form of the oxide of aluminium which occurs as a mineral; it stands second, next to diamond, in the scale of hardness (529). It was formerly much used under the name of emery for polishing hard surfaces, but has been largely replaced by carborundum, SiC, an electric furnace product (218). Ruby and sapphire are crys- talline forms of the oxide, the color being produced probably in the case of the former by a trace of chromium compounds and of the latter by aluminates of iron and titanium. Synthetic rubies and sapphires are now manufactured by fusing in the oxyhydrogen flame aluminium oxide to which has been added the materials required to produce the proper color. Aluminium oxide is often called alumina. Aluminium oxide is used in making chemical apparatus designed to resist high temperatures. The object is fashioned out of the powdered oxide and heated in an electric furnace until the outside has just fused. Although the oxide of aluminium pre- pared by dehydrating the hydroxide is readily soluble in acids, it becomes very inactive after fusion. Ware prepared in this way is sold under the trade name alundum. 570 'INORGANIC CHEMISTRY FOR COLLEGES 687. Aluminium Chloride. When a solution prepared by dissolving aluminium or its oxide or hydroxide in hydrochloric acid is evaporated, the chloride separates in crystals, which have the formula AlCl3,6H2O. The compound is completely hydro- lyzed when an attempt is made to dehydrate it. The anhydrous salt, which is much used in the preparation of certain organic compounds, is prepared by the action of chlorine on aluminium. It has a vapor pressure of 760 mm. at 183, and sublimes without melting, giving a white crystalline solid. The other halides of aluminium are made in a similar way; they all form with the halides of the alkali metals, characteristic double salts of which cryolite, AlFs,3NaF, is a noteworthy example. 688. Aluminium Sulphate. This salt, A1 2 (SO 4 )3,18H 2 O, is extensively used in the industries and is manufactured in large quantities by the action of sulphuric acid on clay, bauxite, or aluminium hydroxide prepared from cryolite. Aluminium sulphate is used in purifying water. It is con- verted, after being dissolved, into the hydroxide by the addition of slaked lime, if the water is soft, or by the calcium bicarbonate present, if the water is temporarily hard. The hydroxide is formed as a gelatinous precipitate that causes the finely divided suspended matter to coagulate into large particles which settle and thus leave the water clear. The precipitate absorbs most of the bac- teria present. Aluminium salts are used in mordanting cotton and linen before they are dyed. These fibers do not absorb from solutions many kinds of dyes in such a form that they are permanently held and not removed by washing. To overcome this, the material to be dyed is first mordanted, that is, an insoluble substance is deposited on the fiber, which unites with the dye and renders it insoluble in water. When aluminium salts are used, the fiber absorbs the aluminium hydroxide formed as the result of their hydrolysis, and the hydroxide unites with the color- ing matter of the d} 7 e-bath. Wool absorbs readily aluminium hydroxide from a solution of the sulphate, but since cotton and linen do not, they are mordanted in a bath containing a soluble basic sulphate formed by dissolving aluminium hydroxide in a solution of the sulphate. Aluminium acetate is used as a mordant ALUMINIUM 571 for cotton with certain dyes; it is the salt of a weak acid and is highly hydrolyzed, and the volatile acid formed is lost by evapora- tion when the cloth dries. Aluminium sulphate is used in making white leather. The hide, after the removal of the hair, is soaked in a bath of aluminium sulphate and salt. The change that takes place is similar to that in the case of mordanting wool. The animal proteins of which the skin is composed absorb aluminium hydroxide and are changed into a form which does not swell in water the skin is, thus, tanned and converted into leather. Chromium sulphate is used in tanning hides, but as its hydroxide is colored it cannot be employed when white leather is desired. Aluminium salts are used in fire-proofing and water-proofing fabrics. Either is accomplished by treating the fabric with alu- minium sulphate and then with a solution of a carbonate, or by wetting it with a solution of aluminium acetate, and then exposing it to steam; the acetate is hydrolyzed, the aluminium hydroxide is deposited in the fiber, and the acid formed passes off with the steam. Alum, K2SC>4,Al2(SO4)3,24H2O, is a member of a class of double sulphates which are composed of 1 molecule of a sulphate of an alkali metal or ammonium, 1 molecule of a sulphate of a trivalent metal, (Al,Cr,Fe,Mn), and 24 molecules of water of crystallization; all the members of this series of compounds are called alums; for example, ferric ammonium alum has the formula (NH 4 )2SO4,Fe 2 (SO4)3,24H 2 O. The salts crystallize as octahedra and are isomorphous. Alum was formerly used for the purposes enumerated above in the case of aluminium sulphate, because it crystallizes well from water and can be readily obtained in a pure condition. It is still used to some extent, but improved methods of preparation of the simple sulphate in the pure condition have done away with the necessity of using the more expensive salt containing potassium. 689. Clay. When feldspar undergoes decomposition as the result of atmospheric influences, it is converted into a hydrated silicate of aluminium, which is called kaolin, if pure, or clay, if it is mixed with other substances. The relation between the compo- sition of feldspar and that of kaolin can be seen from the following formulas : 572 INORGANIC CHEMISTRY FOR COLLEGES Feldspar, KAlSi 3 8 or K 2 0,Al 2 3 ,6SiO 2 Kaolin H 2 Al2(Si04)2,H 2 or 2H 2 O, Al 2 3 ,2Si0 2 The potassium silicate produced in the weathering of the mineral is soluble, and when washed away leaves pure kaolin. It is in this way that soluble potassium salts are formed in the soil. If feld- spar is a constituent of a rock, the other materials remain mixed with the clay and constitute the soil; in the case of granite, the silicon dioxide present is left as sand and the mica is converted into aluminium silicate and potassium silicate. The mineral con- stituents of soils vary with the composition of the rock from which they were formed and the extent to which disintegration of the latter has taken place. The presence of nitrates, fluorides, and phosphates, and of calcium, magnesium, and sodium in soil has already been noted. In addition to these mineral constituents, soils contain a mixture of complex organic substances, known as humus, formed as a result of the disintegration of the vegetable material left in the soil when plants growing on it die. Clays may remain in the place where they were formed, or the kaolin, which is light, may be washed away by streams and be deposited in another place. Such sedimentary clays are usually impure as the result of the admixture of other substances; they are called clays if sufficient kaolin is present in them to form a plastic mass when they are mixed with water. Impure clays of this kind are used in making brick, sewer pipe, tile, and the cheaper varieties of stoneware. The clay is made into a plastic mass with water, molded, dried, and fired in a kiln at about 1000-1200. The substances present in clay other than kaolin have a marked effect on the properties of the finished product. If salts of sodium, potassium, magnesium, calcium, or iron are present, they react to some extent at the temperature of " burning " and form with the kaolin, silicates which fuse, more or less. In this way they bind together the clay particles; as a result, the material after " burn- ing " is partially vitrified and has great strength. If the propor- tion of the metallic compounds is small, the finished product is porous and lacks luster. The color of clay products is due to the presence of iron, which is present in the clay, in all probability, as a colorless hydrated silicate. When the clay is "burned" the salt is decomposed, and ALUMINIUM 573 the iron is oxidized to ferric oxide, Fe20s, which produces a red color. If the clay is burned in a reducing atmosphere ferrous silicate is formed and the color is purplish to black. If iron is absent, the clay burns white, and if present in small quantities, a buff color is produced. The white efflorescence often seen on brick is produced as the result of the solvent action of water on the soluble impurities present in the clay, which have not been converted into silicates when the clay was burned. The deposit may contain the sulphates of sodium, potassium, magnesium, and calcium. The clay products of the cheaper varieties are glazed by introducing salt in the kiln after the material has been fired; at the high temperature, 1200, the kaolin on the surface reacts with the salt, liberates hydro- chloric acid, and is converted into a sodium-aluminium silicate, which fuses. 690. Cast iron is covered with a clay enamel in making bath- tubs, etc., by coating the casting with a mixture prepared from clay, feldspar, sand, and fluorspar or calcite; cryolite is added to render the enamel opaque, and borax makes it more ductile and elastic. The materials are first melted together and then poured into water while in the molten condition. This treatment con- verts the product into a finely divided condition; it is then mixed with more clay and ground with a small amount of water until it takes the form of a thick cream. The iron casting is covered with this material and fired, the process being repeated two or more times. 691. China and porcelain are prepared from pure clay, free from iron, to which has been added more or less feldspar, the amount being determined by the quality of ware desired. It is not necessary in burning the ware containing the larger propor- tion of the more fusible feldspar to heat it to such a high tem- perature, but it is not of such high quality. To the mixture of finely ground material is added just enough water to convert it into a plastic dough-like mass, which is then shaped either on a potter's wheel or in a mold made of plaster of Paris. The material is next allowed to dry and is then fired in a kiln. It is glazed by dipping it into a cream made of feldspar and silica, and then refiring it to melt the feldspar, which flows into the pores of the unglazed material and produces a smooth surface. 574 INORGANIC CHEMISTRY FOR COLLEGES Lead oxide is used in some glazes to lower the temperature of firing, but the product is not so durable; it is not used in making the highest quality china. 692. The plasticity of clay the property which makes it possible to convert it into a dough-like mass with water is due to the colloidal nature of the particles of which it is composed. These are in a very finely divided condition and remain more or less in suspension when mixed with water, on account of the fact that they are charged with negative electricity and repel one another. They do not, therefore, tend to coalesce to larger par- ticles, which would settle out on account of their weight. Clay is rendered colloidal by the presence of certain organic substances produced as the result of the disintegration of the organic matter derived from plants. This material often gets into streams, which are thereby rendered muddy. When these streams enter the sea the salts present cause the precipitation of the colloidal clay, which often forms extensive bars at the mouths of rivers. These facts led to an important invention by Acheson, who first pre- pared artificial graphite in the electric furnace. He showed that if the graphite is ground with tannin or other organic compounds similar to those obtained from plants which render clay colloidal (534), the carbon is converted into a colloidal condition, and when mixed with water or oil makes an excellent lubricant for many purposes. Colloidal clay is removed from water by adding to it aluminium salts, which cause the precipitation of the colloid. A process is now in use for the removal of ferric oxide from clay, which is based on the colloidal properties of these materials. Clay is a negative colloid, being charged with negative electricity, and ferric oxide is a positive colloid. When an electric current is passed through a suspension of the mixture, the clay travels with the negative current and accumulates at the positive pole, and the ferric oxide moves in the opposite direction. The material which settles out at the positive pole is free from iron and is used in making white por- celain. 693. Fuller's earth is a form of clay which is very porous; it is used as an absorbent for colored substances in decolorizing oils. Highly colored clays are used as pigments in making paints; yellow ocher is a clay colored with hydrated oxide of iron, and ALUMINIUM 575 siennas and umbers contain in addition some manganese dioxide. If the material is heated to a high temperature the shade is deep- ened; red ocher is made in this way; burnt sienna is a reddish orange. 694. Portland Cement. The manufacture of cement has become one of the most important chemical industries, on account of the durable properties of materials constructed from cement and the fact that the substances from which it is prepared are abundant. The Romans used for building purposes a rock found near Mount Vesuvius, which had the property of hardening under water. It was a rock of volcanic origin, which was a natural cement. In 1756 Smeaton conducted experiments to determine the best kind of mortar to be used under water in connection with the erection of a lighthouse on Eddystone Rock, and found that one made by burning lime containing clay gave the best results. In 1824 a patent was granted in England to cover the manufacture of a cement made from chalk and mud rich in clay, which had the property of setting under water; it was called Portland cement because when it set, a product was produced which resembled in appearance a limestone quarried near Portland, England. Cement is formed by heating together substances which con- tain the oxides of calcium, aluminium, and silicon. The materials ordinarily used as a source of calcium are limestone, chalk, or marl, which is an amorphous form of calcium carbonate mixed with organic matter and water. The aluminium and silicon are obtained from clay or slate, but in certain cases blast-furnace slag is now used in large quantities. Cement consists of a mixture of calcium silicates and calcium aluminates in the proportions which have been found as the result of practice to give the best results. The materials are mixed in such proportions that the percentage of CaO in the mixture is about twice the sum of the percentages of Si(>2 and A^Os. The relation between the quantities of the acid-forming oxides is approximately 1 AfeOa to 2.5-4 SiCb. The separate materials to be used are crushed, dried, and analyzed, and the amounts of each required to produce a cement of the above composition are mixed. The charge is then ground and passed through a rotary kiln from 70 to 150 feet long and 6 to 8 feet in diameter constructed of steel and lined with fire-brick. 576 INORGANIC CHEMISTRY FOR COLLEGES The kiln is set at an angle of about 15 and rotated about once a minute. As the powdered material tumbles through the kiln it is heated by a flame produced within the kiln by burning gas, oil, or a jet of powdered coal. The material is heated to incipient fusion and forms lumps. The " clinker " so produced is mixed with about 2 per cent of gypsum and ground to a fine powder. The sulphate is added to retard the setting of the cement when it is mixed with water. The chemical composition of cement and the changes which take place when it sets with water have been very fully studied, but no universally accepted explanation of these changes has been reached. It is thought that at least two calcium silicates and two calcium aluminates are present in the cement before it is treated with water; to these have been assigned the following formulas: (CaO) 3 ,Si0 2 , (CaO) 2 ,Si0 2 , (CaO) 3 ,Al 2 3 and (CaO) 2 ,Al 2 O 3 . The setting of cement is due to the changes which take place when water converts one or more of these substances into hydrates and brings about their hydrolysis with the liberation of free calcium hydroxide. In several regions in the United States rocks occur which con- tain their constituents in such proportions that they can be con- verted by heating directly into cement. Such rocks occur in Ohio, Pennsylvania, Illinois, Wisconsin, and Colorado. Large amounts of cement are made from blast-furnace slag (739) and limestone. 695. Tests for Aluminium Salts. Solutions of alkalies, car- bonates, and sulphides precipitate from salts of the metal color- less, gelatinous aluminium hydroxide, which is soluble in sodium hydroxide and insoluble in ammonia. When barium carbonate is shaken with a solution of a salt of aluminium, the metal is pre- cipitated as hydroxide as the result of the fact that the hydrolysis of the salt is made complete owing to the interaction of the car- bonate with the acid formed: A1C1 3 + 3H 2 A1(OH) 3 + 3HC1 BaCO 3 + 2HC1 ^ BaCl 2 + H 2 + C0 2 Salts of the other trivalent metals behave in a similar way, but those of the bivalent metals are not decomposed because they are ALUMINIUM 577 less hydrolyzed by water. The reaction serves, therefore, as a means of separating the trivalent and bivalent metals and is used in analytical chemistry. EXERCISES 1. In the preparation of aluminium by electrolysis, why is sodium not set free from the electrolyte used? 2. The alloys of Al and Cu vary markedly in properties as the proportions of the two metals vary. This is not the case with the Cu-Ag alloys. Can you state a reason for the difference? 3. What compounds could be used in place of the oxides of iron in ther- mite? Could the aluminium be replaced by any other metal? If so what? 4. Starting with Al write equations for reactions by which the following compounds could be prepared: (a) A1(NO 3 ) 3 , (6) alum, (c) NaAlC>2, and (d) A1PO 4 , which is insoluble. 5. Write equations for the reactions which take place between the follow- ing substances in aqueous solution: (a) A1C1 3 and Na 2 S, (b) A1 2 (SO 4 )3 and Na 2 CO 3 , (c) A1(NO 3 ) 3 and NaOH. 6. State as many ways as possible by which the sulphates of the following metals can be distinguished from one another: Ca, Mg, Zn, Al. 7. If you were given a mixture of A1C1 3 and MgCl 2 , how could you pre- pare from it pure MgS0 4 , 7H 2 and pure A1 2 (SO 4 ) 3 ,18H 2 0? CHAPTER XXXIX TIN AND LEAD 696. Tin and lead, which are members of the second family of the fourth group in the periodic classification of the elements, are characteristic of the metals which show the valence 4. The com- pounds derived from them in the bivalent condition resemble, in general, the analogous derivatives of other metals with the valence 2, but then* hydroxides are less basic than those of the metals of the calcium family, and their salts are more highly hydrolyzed as a consequence. When tin and lead exhibit the valence 4 they function as acid-forming elements; stannic chloride, SnCU, for example, is practically not ionized in solution, and is slowly but completely hydrolyzed by water. The fact that increase in valence leads to an increase in acid-forming properties has been repeatedly mentioned; it is clearly illustrated in the case of these two metals. Increase in atomic weight in the family is associated, as is usual in most cases, with increase in base-forming properties. German- ium, the first member of the family, is a weak acid-forming element and shows more resemblances to carbon than to tin. Its monoxide, GeO, is practically neutral like carbon monoxide, and its dioxide, GeO2, is acid-forming. Lead is the most active member of the family in base-forming properties. By considering the above facts in connection with those summarized at the beginning of the last chapter in regard to the general behavior of univalent, bivalent, and trivalent metals, it will be seen that there is a well- defined gradation in chemical properties with increase in valence. When we pass to the next group the fifth the members of which have been described, we find that the basic properties of the ele- ments are exceedingly weak, and that in most cases the important compounds are those in which the elements function as acid- formers. In the sixth group, which contains sulphur, we find a 578 TIN AND LEAD 579 strong acid-forming element, but some of the elements in the group in their lower valencies show weakly metallic properties. TIN 697. Tin was known to the ancients and was used in bronze in prehistoric times. The metal is obtained from cassiterite, which is commonly called tin-stone, SnO2. The ore is first roasted to remove sulphur, arsenic, and antimony, and to oxidize any iron present. It is then treated with hydrochloric acid to convert the impurities in the ore into chlorides, which are washed out with water. The oxide of tin is then reduced with coal. Large quanti- ties of tin-stone were formerly mined at Cornwall, England, but the chief sources of the metal at present are the Malay Peninsula, Bolivia, and the island of Banca in the Indian Ocean. The chief physical and chemical properties of the metal have already been given (table, page 443 and 548). It exists in two forms. One is amorphous, has the specific gravity 5.85, and is stable below 18; the other, which is crystalline and stable above 18, can exist below this temperature, but when it is kept at 15 it passes slowly into the amorphous form. Tin has a low tenacity, but is very malleable and can be made into foil as thin as 0.01 of an inch. It is most malleable at about 100, and becomes so brittle at 200 it can be powdered. It melts at 232, boils at 2270 and has the specific gravity 7.3. The action of the air and acids on tin has been described (548, 550, 549). Tin is used under the name block tin or sheet tin in the manu- facture of pipe and of vessels of various kinds, when a metal is desired which is not attacked by dilute vegetable acids. On account of the high cost of tin other metals, such as copper and iron, are coated with it. Tin-plate is made by coating sheets of steel with molten tin (551) ; it is apt to corrode rapidly when the steel is exposed in any way (567) . Large quantities of tin are used in making alloys, such as solder and bronze (542) . 698. Oxides and Hydroxides of Tin. When stannous salts are treated in solution with bases, stannous hydroxide is formed as a colorless precipitate: SnCl 2 + 2NaOH = Sn(OH) 2 + 2NaCl 580 INORGANIC CHEMISTRY FOR COLLEGES The hydroxide dissolves in an excess of an alkali to form a stannite : Sn(OH) 2 + 2NaOH = Sn(ONa) 2 + 2H 2 O The salt is highly hydrolyzed and is unstable; when the solution is boiled, it is converted into tin and sodium stannate, which resembles sodium carbonate in composition, Na 2 SnO3, and is derived from SnO 2 : 2Sn(ONa) 2 + H 2 O = Sn + Na 2 SnO 3 + 2NaOH Stannous hydroxide is precipitated when a soluble carbonate is added to a stannous salt, the reaction being produced by the same cause that leads to the precipitation of the hydroxides of the trivalent metals by carbonates (682). The base-forming property of stannous tin is exceedingly weak. The fact is of interest in connection with the behavior of salts of other bivalent metals with soluble carbonates; those of calcium yield the normal carbonate, those of zinc, a basic carbonate, and those of stannous tin, the base itself. Stannous oxide, SnO, is obtained by heating stannous oxalate, SnC 2 O4, in the absence of oxygen; it is a black powder, which burns in the air to the white dioxide, Sn0 2 . 699. When bases are added to a solution of a stannic com- pound, SnCU, for example, a white, gelatinous precipitate is formed, which is an unstable hydroxide that loses water gradually until stannic oxide, SnO 2 , is formed. The precipitate dissolves in solutions of alkalies and yields salts of stannic acid, which is derived from the normal hydroxide by loss of water: Sn(OH) 4 - H 2 O = H 2 SnO 3 Stannic acid and its salts are sometimes distinguished from a second stannic acid, formed by the action of concentrated nitric acid on tin, by prefixing to their names the Greek letter alpha, a. 700. When tin is treated with concentrated nitric acid it is converted into a hydrated derivative of the dioxide, which differs in properties from the stannic acid precipitated by alkalies from stannic compounds; it is called metastannic acid or /3-stannic acid. It dissolves with difficulty in boiling solutions of alkalies, from which salts can be obtained; the sodium salt made in this way has the composition represented by the formula TIN AND LEAD 581 For this reason the formula of metastannic acid is generally written (SnO2)5,^*H20; the amount of water in combination varies with the conditions under which the acid is made. When the acid is fused with sodium hydroxide, it is converted into sodium stannate, Na2SnOs, and when heated alone, into the dioxide. Tin dioxide is white when cold, but resembles zinc oxide in turning yellow when hot. The oxide that occurs as cassiterite is black, the color being due to the presence of oxides of iron. Tin dioxide is used in fire-proofing fabrics. The materials are first soaked in a solution of sodium stannate and then treated with a solution of ammonium sulphate. The reaction which takes place leads to the precipitation in the fiber of stannic acid, which loses water and passes into the dioxide. 701. The Chlorides of Tin. Stannous chloride is formed when tin dissolves in hydrochloric acid. The salt obtained from the solution by crystallization has the formula SnCl2,2H2O. It is highly hydrolyzed, and when a strong solution of it is poured into water a precipitate of a basic salt, Sn(OH)Cl, is formed. Stannous chloride is oxidized by the air and is slowly converted in solution into stannic chloride, SnCU, and basic stannous chloride, Sn(OH)Cl, which precipitates. The salt is kept in the stannous condition and the precipitation of the basic salt avoided by adding to the solution a small amount of hydrochloric acid and some metallic tin The tendency of stannous tin to pass into stannic tin is also shown when it is brought into contact with salts that can be reduced. Ferric chloride, for example, is reduced to ferrous chlo- ride 2FeCl 3 + SnCl 2 = 2FeCl 2 + SnCU and mercuric chloride, to mercurous chloride and mercury (674). Stannous chloride is used as a mordant under the name tin crystals. Stannic chloride is made from tin and chlorine. The anhy- drous compound is a colorless, fuming liquid, which boils at 114 and dissolves in water. It forms a number of hydrates, of which the one having the formula SnCl4,5H20 is the most important. It is used as a mordant on account of the fact that it is appreciably hydrolyzed and the hydroxide formed is absorbed by the fiber and holds the dye. Large quantities of it are used in making the cheaper grades of silk, because it adds materially to the weight of the latter and improves its appearance. Silks weighted in this 582 INORGANIC CHEMISTRY FOR COLLEGES way contain at times as much as 75 per cent of the oxide and are not durable. " Pink salt," which is extensively used as a mordant, is a double salt of the formula (NH^SnCle. A large number of double salts containing stannic tin are known; their easy forma- tion is traceable to the acid-forming properties of the metal. 702. The Sulphides of Tin. Stannous sulphide, SnS, obtained by the action of hydrogen sulphide on a solution of stannous chloride 1 is a brown precipitate, which is insoluble in ammonium sulphide but dissolves in ammonium polysulphide as the result of the formation of ammonium thiostannate, (NH^SnSa. Addition of an acid causes the precipitation of yellow stannic sulphide, SnS2, which is also formed from stannic chloride and hydrogen sulphide. The reactions are analogous to those dis- cussed at length in the case of arsenic and antimony (478) and, as a result, in qualitative analysis tin falls into the group which includes these elements. 703. Test for Tin Salts. The properties of the sulphides just given are utilized in analytical chemistry. A confirmatory test lor tin is based on the fact that metallic zinc sets free tin from its salts, and the metal obtained in this way is soluble in dilute hydro- chloric acid. LEAD 704. Metallurgy of Lead. Lead has been known since pre- historic times. It occurs as the carbonate and the sulphate, which are used as ores of the metal, but the most important source is the sulphide, PbS, galena, which crystallizes in black cubes having a high silvery luster. The metal is ordinarily obtained from the ore by first heating it in a stream of air to convert it into the oxide, and then reducing the latter in a blast furnace with coal. Many ores of lead contain silver and gold, and a part of the metallurgical processes used in extracting lead have to do with the separation of the noble metals; these will be considered later (713). 705. Properties of Lead. The more important physical prop- erties of lead have already been given (table, page 443). Lead has a bluish-gray color with a marked metallic luster when freshly cut, but it soon tarnishes in the air as the result of the formation of a thin coating of oxide, which changes to a basic carbonate. TIN AND LEAD 583 The coating is closely adherent and retards further corrosion to such an extent that lead is used when a metal is required that is stable under atmospheric influences. The metal is very inactive toward acids, and is used in the manufacture of many kinds of apparatus designed for the chemical industries. Its use in alloys has already been explained (542). Lead is used in large quantities in the manufacture of pipe, which is made by forcing the metal while soft through dies. 706. Oxides of Lead. When the metal is heated in the air it is converted at a temperature just above its melting-point, 327, into a yellow monoxide, PbO, which is called massicot. If the temperature is such that the oxide melts, the product contains a trace of 'red lead and is reddish yellow; it is known as litharge. Red lead, PbsO 4 , is prepared by first converting lead into massi- cot and then heating it at about 480 in air, the process requiring about forty-eight hours. This oxide is used in glass making and in paints to protect iron and steel from rusting. It reacts with linseed oil to form a hard solid, and for this reason a mixture of the two substances is used as a lute in plumbing and gas fitting. A mixture of red lead and glycerine is used as a cement, because on standing it is converted into a hard, adhering mass. When red lead is heated to a high temperature it gives off oxygen and is converted into the monoxide. Warm dilute nitric acid converts it into lead nitrate, Pb(NOs)2, which dissolves, and lead dioxide, PbO2, which is left as a brown powder. For this reason red lead is thought to be a salt of the formula Pb 2 (Pb0 4 ), which is derived from a hydroxide of the composition Pb(OH)4 by replacing the hydrogen atoms by two atoms of bivalent lead. When lead has the valence 2 it has base-forming properties; in the quadrivalent condition it is acidic and its hydroxide, Pb(OH)4, is called orthoplumbic acid. From this point of view the reaction between nitric acid and red lead, which is lead orthoplumbate, is written as follows : Pb 2 Pb0 4 + 4HN0 3 = 2Pb(N0 3 ) 2 + H 4 Pb0 4 H 4 PbO 4 = Pb0 2 + 2H 2 O Lead dioxide , Pb0 2 , is a brown powder, insoluble in water, which is usually prepared by treating with bleaching powder a 584 INORGANIC CHEMISTRY FOR COLLEGES solution of sodium plumbite, Na2pb02, formed by dissolving lead hydroxide in caustic soda. The sodium salt hydrolyzes giving sodium hydroxide and lead hydroxide, which is oxidized by the oxygen furnished by the bleaching powder : Na 2 PbO2 + CaOCl 2 + H 2 O = 2NaOH + CaCl 2 + PbO 2 . The dioxide dissolves in strong solutions of alkalies, playing the part of an acid anhydride, and forms salts related in composition to the stannates and carbonates. Potassium plumbate is obtained from aqueous solutions in crystals, which have the formula K2PbO3,3H2O. Lead dioxide reacts with hydrochloric acid in a way entirely analogous to that which has been described in the case of man- ganese dioxide (112): PbO 2 + 4HC1 = PbCl 2 + 2H 2 O + C1 2 . The use of lead dioxide in storage batteries has been described at some length in section 578. Lead hydroxide, Pb(OH) 2 , is formed as a white precipitate when a solution of a base is added to a lead salt; it reacts with caustic alkalies and forms soluble plumbites, but does not dissolve in ammonia. 707. The Chlorides of Lead. Lead chloride, PbCl 2 , is diffi- cultly soluble in cold water and is precipitated when a soluble chloride is added to a solution of a lead salt ; it is colorless and dis- solves in hot water, from which it crystallizes when the solution cools. Lead tetrachloride, PbCU, is stable only at low temperatures; it is a liquid which fumes in the air and decomposes with a small amount of water to form the dichloride and chlorine. With larger amounts of water it is hydrolyzed and converted into lead dioxide and hydrochloric acid. Double salts containing lead tetrachloride are comparatively stable, and advantage is taken of this fact in preparing the compound. When chlorine is passed into a mixture of lead chloride and hydrochloric acid, the lead tetrachloride formed unites with some of the acid and forms a complex acid of the composition H 2 PbCle (PbCU,2HCl). Ammonium chloride is next added to the solution to form ammonium chloroplumbate, (NH4) 2 PbCl6. When this salt is added to concentrated sulphuric acid at a low temperature, the salt is decomposed and lead tetra- chloride separates as a heavy oil. Lead iodide, PbI 2 , has characteristic physical properties and is useful in identifying lead compounds. It is precipitated as a TIN AND LEAD 585 yellow powder when a soluble iodide is added to a lead salt. It crystallizes from boiling water in yellow scales, which have a brilliant luster. 708. Lead Carbonate. The neutral carbonate of lead, PbCOs, occurs as the mineral cerussite, and was used by the Romans as a white pigment. It was early replaced by a basic carbonate for this purpose, and a method of preparing the latter used in Hol- land in the seventeenth century is still employed in the manu- facture of " white lead," which is to-day the most important pigment used in making paint. The value of a pigment is largely determined by its so-called " covering power," which depends on its opacity and physical condition. White lead as ordinarily made has a composition closely approaching that of the formula: Pb 3 (OH) 2 (CO 3 )2 or Pb(OH) 2 ,2PbCO 3 . A compound of the same formula is obtained when sodium carbonate is added to a soluble lead salt, but it does not have the same covering power as white lead. Many attempts have been made to replace the old Dutch process for making the pigment, for it requires a long time. Some of these have been successful, but the white lead prepared by the older process is held by some to have superior qualities. In the Dutch process for making white lead, sheet lead or cast- ings made from the metal in such a form that they have a large surface, are placed on shelves in earthenware pots, the bottoms of which are covered with a dilute solution of acetic acid, HC2H3O2. The pots are set on a layer of moist spent tan bark. A floor is next placed a few inches above the tops of the open pots, and this is covered with tan bark, upon which more pots are set. The process is repeated until the room is filled. It is then closed and left for about three months. The bark soon begins to ferment and, as the result, carbon dioxide is formed and the temperature rises to 55-60. Loss of heat is prevented by having a thick layer of the bark around the walls of the room. Acetic acid is volatilized and slowly converts the lead into an acetate, which is changed by the water and carbon dioxide into the basic carbonate, the acetic acid being regenerated. When the corroded material is taken from the pots it is found to be in the form of a hard com- pact mass, which retains the original shape of the lead. It is ground to a fine powder under water, and the milky liquid is allowed to settle. After the larger particles have separated, the liquid is 586 INORGANIC CHEMISTRY FOR COLLEGES withdrawn, and on further settling the white lead is obtained as a heavy mud. This is then mixed with a little linseed oil, which causes the separation of the water. The mixture of white lead and oil is separated and packed. This method of treating the mud avoids handling dry, powdered material, which is poisonous and is apt to get into the air as dust. In Carter's process, which is much used, lead is first converted into a very finely divided condition by blowing a jet of super- heated steam against a stream of molten lead. The " atomized " lead is then tumbled in a rotating cylinder into which acetic acid and carbon dioxide are led. The corrosion is complete in about fifteen days. White lead is an excellent pigment and is the basis for most colored paints, which are prepared by the addition of colored substances to a mixture of white lead and linseed oil. White lead is converted by hydrogen sulphide into lead sulphide, which is black; it cannot be used in paints which are exposed to this gas. Under these circumstances zinc white, which forms a white sul- phide, is used. 709. Other Salts of Lead. Lead nitrate, Pb(NO 3 ) 2 , is the common laboratory reagent. It is made by dissolving lead oxide . in nitric acid; it crystallizes readily from hot water containing a small amount of nitric acid, which prevents its hydrolysis. Lead sulphate, PbSO*, is a very insoluble salt and is made by precipitation; it dissolves, like barium sulphate, in concentrated sulphuric acid. It is soluble in concentrated solutions of the alkalies as the result of the formation of sodium plumbite. Lead acetate, Pb^HsC^, called sugar of lead, is prepared by the action of air and acetic acid on lead, or by dissolving lead oxide in the acid. It is used in the preparation of a basic lead acetate, which is formed when a solution of the neutral salt is boiled with litharge. Basic lead acetate has the formula Pb(OH)C2H3O2; it is soluble in water and is used as a mordant. Lead sulphide, PbS, is formed as a black precipitate from lead salts and soluble sulphides. It is insoluble in dilute hydrochloric acid, but dissolves in concentrated nitric acid. 710. Tests for Lead Salts. The formation of a black sulphide in the presence of dilute acids when a lead salt is treated with hydrogen sulphide, and the formation of a white insoluble sulphate TIN AND LEAD 587 when lead salts are treated with a soluble sulphate, are used as a test for the metal. Sodium hydroxide precipitates white lead hydroxide which is soluble in an excess of the reagent but insoluble in ammonium hydroxide. Lead appears in the first group in qual- itative analysis as it is precipitated as chloride when hydrochloric acid is added to a solution containing it. It is separated from the insoluble chlorides by dissolving it from the mixture with boiling water, from which it separates on cooling. Since lead chloride is slightly soluble in water (1.5 parts in 100 at 18), lead is not com- pletely removed from solution when it is precipitated as chloride. EXERCISES 1. Summarize in the form of a table the properties of metallic elements and their hydroxides and salts, which are markedly influenced by the valence of the metal. Consider metals having the valence 1, 2, 3, and 4. 2. Write equations for reactions which take place between the following: (a) Sn and concentrated HNO 3 , (6) Sn and dilute HNO ? , (c) SnCl 2 , H 2 O, and air, (d) Na 2 SnO 3 and (NH 4 ) 2 SO 4 . 3. By what chemical properties could you distinguish (a) Sb 2 S 3 from SnS 2 , (6) SnCl 2 from A1C1 3 , (c) pure tin foil from a foil containing Sn and Pb, (d) pure tin. from an alloy of Sn and Zn? 4. How could you determine (a) whether a sample of silk had been mor- danted with tin salt and (6) the percentage of SnO 2 present? 5. How could you determine whether a piece of cotton had been fire- proofed with alum or sodium stannate? 6. Write equations for reactions which would serve to distinguish from one another the following: (a) PbO, (6) PbO 2 , and (c) Pb 3 O 4 . 7. How could you obtain (a) Pb 3 O 4 from a mixture of Pb 3 O 4 and ZnO, (6) MgO from a mixture of MgO and PbO? 8. How could Pb and Cu be separated? 9. (a) Why does lead form double salts readily? What substance could be used to dissolve lead iodide? 10. (a) State a number of ways of determining whether a white paint is made from white lead or zinc white. (6) If the paint contains both, how could you show the presence of each? 11. What substance could be used to restore the white color to a lead paint which has turned yellow in the air? CHAPTER XL COPPER, SILVER, AND GOLD 711. The position of copper, silver, and gold in the first group of the periodic classification is anomalous, because the metals show no marked resemblances to those of the alkalies, and because copper and gold form compounds in which the metals have a greater valence than 1. It is for this reason that their consider- ation has been delayed to this point. We have seen that there is a gradual change in the properties of the compounds derived from metals as the valence of the latter changes from 1 to 4; and we have learned the general properties of the compounds of the four types of metals. With these facts in mind it will be easier to under- stand the chemical behavior of copper, which functions as a uni- valent and bivalent metal, and that of gold, which has the valencies 1 and 3. Silver is always univalent and the properties of its com- pounds are typical of those of other compounds derived from heavy metals when they exhibit this valence. For this reason silver and its compounds will be first considered. SILVER 712. Silver is a relatively strong base-forming element; its hydroxide, which exists only in solution and is formed when silver oxide dissolves in water, is highly dissociated and forms salts which are not hydrolyzed. In this respect it resembles the alkali metals. It is a very inactive element, however; it does not liberate hydrogen from acids, and its oxide is readily decomposed by heat. The metal has been known since the earliest times and was studied by the alchemists, who used a crescent as its symbol in their writings. The symbol used to-day, Ag, is derived from the Latin name of the metal (argentum). 588 COPPER, SILVER, AND GOLD 589 713. Occurrence and Metallurgy of Silver. The metal occurs in the free condition often alloyed with gold and with copper. It occurs as the sulphide, Ag2S, argentite, which is asso- ciated with lead sulphide. The metal is obtained in large quan- tities as a by-product in the smelting of lead and copper. It also occurs as double sulphides with antimony and arsenic, and as silver chloride. Native silver is obtained by amalgamation ; the crushed ore is allowed to stand in contact with mercury, which dissolves the metal, and the solution formed is then distilled; the residue is silver alloyed with any gold present in the ore. The method most commonly used when the ore is a compound of the metal, is to extract it, after pulverization, with a dilute solution of sodium cyanide. The silver is converted into silver cyanide, which dis- solves in an excess of the alkali cyanide to form the complex salt, NaAg(CN)2. Zinc is then added to the solution to precipitate the metal. In the case of some ores, the cyanide and amalgamation processes are used simultaneously. Low-grade silver ores are often smelted with lead ores, since the latter frequently contain silver, which must be recovered. The method commonly employed is called the Parkes process. Silver and gold are much more soluble in zinc than in lead; and since zinc is practically insoluble in lead it can be used to extract the precious metals from the latter. In carrying out the process zinc is well stirred into the molten lead, which is then allowed to cool slowly. A scum consisting of zinc, silver, gold, and some lead separates on the surface of the metal; this is removed and dis- tilled and thus freed from zinc. The residue is then " cupelled," that is, it is heated in a stream of air, which oxidizes the lead and blows the molten oxide to the edge of the hearth, where it is skimmed off. Processes which are used in Mexico and South America are based on other reactions. In one of these the ground ore is mixed with water, salt, and copper sulphate, and heated for several hours. Metallic iron is then added to set silver free from the chloride, and the metal is extracted from the mass by amalgamation. 714. Properties and Uses of Silver. The chief physical prop- erties of silver are given in the table on page 443 and discussed in section 548. The metal is the best conductor of electricity and 590 INORGANIC CHEMISTRY FOR COLLEGES heat; it is exceedingly malleable and ductile. The molten metal possesses the property of dissolving oxygen (about 22 volumes), which is expelled when it solidifies. In the massive or crystalline condition silver is white; when precipitated from solution by zinc it forms a black powder, but mercury causes the separation of the metal from its salts in a highly crystalline condition which leads to the formation of what is called a silver " tree." By the use of reducing agents silver can be obtained in the colloidal condition, the color of the metal varying with the substance used. Silver is a relatively soft metal and for this reason an alloy of the metal is used in coins. Those of the United States and the countries of continental Europe contain 900 parts of silver and 100 parts of copper; the coins of Great Britain and sterling silver contain 925 parts of silver and 75 parts of copper. So-called oxidized silver is made by treating the metal with a solution of a soluble sulphide, which causes the precipitation of silver sul- phide on the surface of the metal. In making mirrors silver is deposited on glass by covering the well-cleaned surface with a solution of a silver salt to which has been added a reducing agent, such as glucose or an ammoniacal solution of cream of tartar. Silver is a very inactive element and does not tarnish in the air except in the presence of hydrogen sulphide. It turns yellow or black when in contact with certain organic products which contain sulphur, such as the protein in eggs. The behavior of the metal with acids has been discussed. It dissolves in nitric and concentrated sulphuric acids; it is not attacked by alkalies, even when they are fused, because it does not play the part of an acid- forming element. Large quantities of silver are used in plating other metals (583) and in the preparation of silver nitrate, which is the source of the compounds of the metal used in photography. 715. Oxides of Silver. When a base is added to a solution of a silver salt a brown precipitate of silver oxide, Ag2O, is formed. The oxide is slightly soluble in water, and the solution contains silver hydroxide, which has not been isolated; it is a relatively strong base and yields salts which are not hydrolyzed by water. The oxide dissolves in ammonia and forms a compound of the com- position Ag(NH3)2OH, which is converted slowly, but more rapidly COPPER, SILVER, AND GOLD 591 in the presence of caustic alkalies, into a black, highly explosive powder, the composition of which is unknown. Silver peroxide, Ag2C>2, is formed as the result of the action of ozone on silver. 716. The Halides of Silver. The chloride, bromide, and iodide of silver are very slightly soluble salts and are obtained as curdy precipitates when the respective ions are added to a solution of a silver salt. The chloride is white, the bromide pale yellow, and the iodide a deeper yellow; the solubilities of the three salts are, respectively, O.OslS, O.C^l, and 0.0e35 gram in 100 grams of water at 18. Silver chloride dissolves readily and silver bromide less readily in a solution of ammonia and form salts of the composition Ag(NH 3 ) 2 Cl and Ag(NH 3 ) 2 Br. The insoluble salts of silver dis- solve in potassium cyanide and form a soluble complex cyanide of the formula KAg(CN)2. Active metals like zinc displace silver from all its soluble salts. Moist silver chloride rapidly changes in color from white to violet when exposed to the sunlight. It is probable that the color is due to the presence in the salt of a small amount of colloidal silver, which is produced as the result of the conversion of a part of the chloride under the influence of light into silver and free chlorine, which escapes. The halides of gold, mercury, lead, and other metals are also reduced in the presence of sunlight. 717. Silver Nitrate. This salt is prepared by dissolving silver in nitric acid. It forms colorless anhydrous crystals, which melt at 208.6. Like the nitrates of the alkali metals, it is more stable when heated than the corresponding salts of the bivalent metals. Advantage is taken of this fact in purifying silver nitrate pre- pared from the metal which contains copper. When the nitrates are heated the copper salt is converted into copper oxide, oxygen, and oxides of nitrogen; the silver salt is not decomposed, and is extracted from the mass after cooling. Silver nitrate, which has been fused and cast into sticks, is used under the name lunar caustic to cauterize sores. The salt is also used in some forms of ink to mark fabrics. When it is in contact with organic material, such as cotton or wool, and is heated or exposed to light, it is reduced and metallic silver is deposited. Inks of this kind should not be used in marking clothing that is to be washed in laundries, because the silver is 592 INORGANIC CHEMISTRY FOR COLLEGES converted into the colorless chloride by the bleaching solution commonly used. 718. Test for Silver Salts. The properties of the halides of silver (716) are utilized in testing for the metal. It is precipitated as chloride in the first group in qualitative analysis when hydro- chloric acid is added to the solution to be analyzed. The chloride is separated from lead chloride by dissolving the latter in boiling water, and from mercurous chloride by treatment with ammonia, which dissolves silver chloride. It is obtained from the solution made in this way by adding nitric acid. Black silver sulphide is precipitated when solutions of silver salts are treated with hydro- gen sulphide. The double cyanides containing silver are not precipitated in the presence of an excess of an alkali cyanide. Silver sulphide is reduced by nascent hydrogen to silver. 719. Photography. The first successful photographic process, which was discovered by Daguerre, was used until the more modern methods were developed. Daguerrotypes were made by exposing in a camera a plate of silvered copper upon the surface of which silver iodide had been deposited by exposing the plate to the vapor of iodine. After exposure, the plate was set over a vessel containing mercury, which was gently heated. The vapor of the metal condensed on the parts of the plate that had been exposed to light and, as a result, the picture was developed. In modern photography the material of the plate or film which is sensitive to light, is prepared by making a jelly from gelatine, in which is precipitated silver bromide by the action of silver nitrate on ammonium bromide. The jelly, or emulsion, as it is called, is first solidified and converted into small pieces, which are washed thoroughly with cold water to remove the soluble salts present. It is then melted and heated for some time in order to obtain the colloidal particles of the silver bromide of such a size that they have the desired sensitiveness to light. After this "ripening" process, the melted emulsion is poured on plates of glass or celluloid films and allowed to cool. A great deal of experim entation has been devoted to a study of the change that takes place when a photographic plate is exposed to light. The problem is a difficult one since the change occurs to such a small extent that its effect cannot be seen or discovered in any way except by "development." It is believed at present that light causes the decomposition of a very small amount of the silver bromide into silver and bromine, which is absorbed by the gel- atin. When the plate is treated with a developer, the silver bromide which is in contact with the trace of metallic silver set free by the light, is reduced to the metal, which, in turn, makes it possible to reduce more bromide. Silver thus serves as a catalytic agent in the reduction of the bromide, and substances are used as developers which reduce silver bromide very slowly in the absence of the metal. In the early days of photography ferrous salts COPPER, SILVER, AND GOLD 593 were used as developers because they reduced silver salts to the metal as the result of their oxidation to ferric salts. At present organic compounds are employed almost exclusively; a solution of pyrogallic acid in an alkali is commonly used as a developer, as it can be readily oxidized. If a plate has been overexposed, it is reduced very rapidly in the developer, and an unsatisfactory negative is produced. To retard development and thus allow the deposition of silver in such a way that normal contrasts between the high lights and shadows are produced, a trace of potassium bromide is added to the developer; this reduces the solubility of silver bromide in water (596) and, consequently, the rate at which silver is deposited from the salt. The plate is next "fixed" by placing it in a bath containing sodium thio- sulphate, commonly called "hypo," which dissolves from the plate the un- reduced silver bromide. The reactions which take place are as follows: 2AgBr+ Na 2 S 2 O 3 = Ag 2 S 2 O 3 + 2NaBr Ag 2 S 2 3 + SNa&O, = 2Na 3 Ag(S 2 O 3 )2 Ordinary alum (688), or better, chromium alum, K 2 S04,Cr 2 (SO 4 )3, 24H 2 O, is at times added to the fixing bath, and is of especial value in warm weather, as it reacts with the gelatin and hardens it so that it holds but a small amount of water, and can be dried more rapidly and with less danger of being scratched. The plate is next washed carefully with water to remove all of the thio- sulphates. because if they remain on the plate they are converted on stand- ing into silver sulphide, which forms a yellow stain. A positive may be made from the negative by printing on a "gaslight" paper which has been coated with an emulsion resembling in composition that used on the plate; it usually contains some silver chloride, and is much less "rapid." After exposure the paper is developed and fixed by the same method used with plates. Some forms of paper are coated with silver chloride, and are so slow in action that they can be handled in daylight. The print is made by exposing the paper under the negative to strong sunlight until the picture is produced. It is then "toned" by placing it in a bath containing a gold salt, NaAuCl 4 ; the silver produced as the result of the action of light on the chloride, passes into solution and gold is deposited in its place. The rate at which the latter is formed determines its color, which may vary from brown to deep purple. The desired shade is obtained by adding to the toning bath substances, such as borax, which affect the rate at which the metal is deposited and the size of the colloidal particles formed. Silver prints can also be toned with platinum. The so-called platinum papers are made with- out the use of silver salts. Ferric oxalate, Fe 2 (C 2 O 4 )3, is the sensitive salt used. When light penetrates the negative it is reduced to ferrous oxalate, FeC 2 Od; and when this is brought into contact with potassium chloroplatinite, K 2 PtCl 4 , the latter is reduced and platinum is deposited. Blue-print paper is coated with a ferric salt as its light-sensitive ingredient and with potassium fer- ricyanide, which reacts with the ferrous salt formed to produce a blue insoluble compound; ferric-ammonium citrate is commonly used, because it dissolves readilv in water and is more sensitive than most other ferric salts. 594 INORGANIC CHEMISTRY FOR COLLEGES The salts of silver are much more rapidly affected by the shorter wave- lengths of light, and, as a consequence, a print from a negative made on an ordinary plate does not show the gradation in brightness of the original object. Bright shades of red and green appear dark, and deep shades of blue appear light. This effect is, in part, overcome by adding to the emulsion used on the plate, dyes that absorb the longer wave-lengths, which then become active. Plates made in this way are called orthochromatic. COPPER 720. Copper was used in prehistoric times for making weapons and tools, and later was alloyed with tin to form bronze, which was the most important metal of the Greeks and Romans. It was replaced for these purposes by iron and steel, but brass, the alloy of the metal with zinc, proved to have properties which adapted it to many uses, and large amounts of copper are used in the manufacture of the alloy. The great development of the electrical industries has resulted in such extensive uses of the metal that it now ranks next to iron in importance. In the stable and more important compounds of copper the metal has the valence 2, but it also forms a few compounds derived from univalent copper. It resembles mercury in this respect, but the series of compounds derived from univalent copper is less com- plete than that of the mercury compounds; no sulphate or nitrate of cuprous copper has been isolated. Copper does not show acid- forming properties as zinc does, but resembles the latter in uniting with ammonia to form complex positive ions, for example, Cu(NH 3 )4 ++ . The salts of the metal are slightly hydroly zed in solution and, like zinc salts, yield a basic carbonate when treated with a soluble carbonate. 721. Occurrence and Metallurgy of Copper. Copper occurs in a large number of minerals, the more important of which are chalcocite, Cu2S, chalcopyrite, CuFeS 2 , bornite, CusFeSs, cuprite, Cu 2 O, azurite, Cu 3 (OHJ 2 (CO 3 )2, and malachite, Cu2(OH) 2 CO 3 . The metal also occurs in the free condition in large quantities in the Lake Superior mines in Michigan. The chief ores mined in Montana are sulphides. On account of the increasing demand for copper, methods have been designed to use as ores, materials which contain as low as 2 per cent of copper compounds mixed with rock. Such ores are COPPER, SILVER, AND GOLD 595 pulverized and " concentrated," by washing them with water down inclined planes; the "gangue," being lighter than the ore, is washed away faster, and a material containing a high precentage of the latter is obtained. This method has been largely super- seded by the oil flotation process, in which the pulverized ore is agitated with water in the presence of a small amount of oil. The latter adheres to the ore but not to the gangue, and, as a conse- quence, the scum which rises to the surface contains the ore, which can be readily separated. Oil flotation is used in concen- trating ores in general, and has made it possible to use low- grade material which could not be economically concentrated by other methods. The methods commonly used in extracting copper from its ores have been described briefly in the chapter in which metallurgy was considered; the electrolytic refining of copper has also been discussed (582). 722. Properties of Copper. The chief physical (table, page 443) and chemical (549, 551) properties of copper as well as its use in alloys (542) and for other purposes have been considered in some detail in previous chapters. The metal becomes coated in the air with a green layer of a basic carbonate and is attacked slowly by non-oxidizing acids if oxygen is present. A deposit of a basic chloride is formed on the metal when it is left in contact with sea- water. 723. The Oxides and Hydroxides of Copper. Cuprous oxide, Cu2O, occurs in nature as cuprite, and on account of its red color is called ruby copper. It is made by reducing an alkaline solution of a copper salt with glucose. The precipitated, hydrated oxide, which is yellowish red, loses most of its water quickly and yields the anhydrous red oxide, Cu2O, when gently heated. Like silver oxide, Ag2O, it dissolves in ammonia; as a result it is converted into the compound of the formula Cu(NHs)2OH, which is colorless. The solution turns blue rapidly if air is present, as the result of oxidation to the cupric compound, Cu(NH3)4(OH)2, which is deep blue. The oxide dissolves in concentrated hydrochloric acid and a compound of cuprous chloride and the acid is formed H^CuCla. Cuprous salts of the oxygen acids are unknown, and when such acids are added to cuprous oxide, a cupric salt and free copper are produced. 596 INORGANIC CHEMISTRY FOR COLLEGES Cupric oxide, CuO, is formed when copper is heated in the air, or when the nitrate, sulphate, carbonate, or hydroxide of the metal is ignited. It is an active oxidizing agent when heated, and is used as such in the analysis of compounds containing carbon and hydrogen. The material to be analyzed is heated in a stream of oxygen in a hard glass tube, and the vapors produced are passed over red-hot copper oxide, which insures their complete combustion. The water formed is collected in a tube containing calcium chloride, and the carbon dioxide, in a bulb containing a solution of potassium hydroxide. 724. Copper hydroxide, Cu(OH)2, is formed as a light-blue precipitate when a solution of a base other than ammonium hydroxide, is added to a copper salt. If the solution contains a large excess of alkali and it is heated, the hydroxide loses a part of its hydrogen and oxygen as water and is converted into a black compound. Cupric hydroxide dissolves in ammonia and forms a deep-blue solution, which contains a compound of the formula Cu(NH3)4(OH)2. Cotton, paper, and other forms of cellulose dis- solve in a strong solution of copper hydroxide in ammonia. The resulting solution is used in making one form of artificial silk by ' forcing it through capillary tubes into a dilute solution of sulphuric acid, which precipitates the cellulose in the form of a fine fiber. Other kinds of artificial silk are made in a similar way by using a solution prepared by the action of sodium hydroxide and carbon disulphide on cellulose (viscose silk), or a solution of acetate of cellulose. The products made in this way are not silk, which is a protein and contains nitrogen, but cellulose in such a form that is has a high luster. Silk dissolves readily in a warm solution of sodium hydroxide, but artificial silk does not. 725. Copper hydroxide dissolves in Rochelle salt, which is sodium-potassium tartrate, NaK(C 4 H 4 O 6 ). The blue solution prepared in this way is called Fehling's solution and is used in the analysis of certain sugars, of which glucose is an example. These compounds when gently heated with Fehling's solution reduce the copper to hydrated cuprous oxide, which is precipitated. The reaction is utilized in testing urine for glucose, which is pro- duced in the body and is excreted in the case of the disease known as diabetes, COPPER, SILVER, AND GOLD 597 726. The Halides of Copper. Cupric chloride crystallizes from water as a blue hydrate, CuCl2,2H2O. When its solution is treated with ammonia, the green basic chloride, CuCl2,3Cu(OH) 2 , first precipitated, dissolves and forms a deep-blue solution from which a compound of the formula Cu(NH3)4d2,H2O can be isolated. Cuprous chloride, CuCl, is prepared by heating a solution of cupric chloride with copper in the presence of strong hydro- chloric acid, and pouring the resulting solution into water. The cuprous chloride first formed CuCk + Cu = 2CuCl dissolves in the excess of the acid and forms a soluble compound of the formula H^CuCls, which decomposes when the solution is diluted: H^CuCls = CuCl + 2HC1. Cuprous chloride is white, insoluble in water, and resembles in some respects silver chloride in its chemical properties. It is soluble in ammonium hydroxide and forms a colorless compound of the formula Cu(NH3)2Cl. Moist cuprous chloride turns violet in the sunlight as the result of reduction. It is oxidized slowly in moist air to a green basic cupric chloride. Ammoniacal solutions of cuprous chloride or other cuprous salts dissolve carbon monoxide, and unstable addition-products are formed, the composition of which is determined by the pressure of the gas in the solution. The reaction is used in the quantitative analysis of flue gases and in the purification of the nitrogen and hydrogen used in the synthetic production of ammonia. Cupric bromide, CuBr2, which is black in the solid condition, forms concentrated solutions which are brown but change on dilu- tion to the blue color characteristic of solutions of copper salts. The change in color is probably due to the formation of copper ions from the undissociated salt. 'Cupric iodide does not exist at ordinary temperatures; when a soluble iodide is added to a solution of a cupric salt, white cuprous iodide is precipitated and iodine is set free: 2CuS0 4 + 4KI = 2CuI + I 2 + 2K 2 SO 4 The reaction is used in the quantitative volumetric analysis of copper compounds; the iodine set free when an iodide is added to the solution is determined by titration with a solution of sodium thiosulphate (310). 598 INORGANIC CHEMISTRY FOR COLLEGES A similar reaction takes place when a solution of potassium cyanide is warmed with a copper salt; the cupric cyanide first formed decomposes into cuprous cyanide, CuCN, and cyanogen, (CN)2, which is a gas. Cuprous cyanide dissolves in potassium cyanide and forms the complex salt KCu(CN)2, which is very stable. 727. Copper Sulphate. This salt, which is called blue vitriol or bluestone, is obtained as a by-product in the " parting " of gold and silver, and also by allowing dilute sulphuric acid to drip on copper scrap in the presence of air. In another process the scrap metal is heated red hot in a furnace and sulphur is added to form copper sulphide; air is next admitted to convert the latter into a mixture of oxide and sulphate, which, on treatment with sul- phuric acid and water and subsequent evaporation yields blue vitriol, CuSO4,5H2O. In very dry air the crystals effloresce and change to a white powder, which has the composition CuSCU,- H2O. This compound, which is prepared by heating the hydrated salt at about 100, is frequently used in the laboratory for drying liquids which react with calcium chloride or potassium hydroxide. Copper sulphate is used in copper plating, as a mordant in calico printing and dyeing, in the preparation of pigments con- taining copper, as a fungicide, and in making germicides and in- secticides. Bordeaux mixture, which is prepared by mixing milk of lime and a solution of copper sulphate, is used to spray plants to prevent the growth of fungi. 728. Other Compounds of Copper. Copper nitrate crystallizes from cold water in deliquescent crystals, Cu(N03)2,6H2O, which are converted into a basic salt when heated at 100. The basic copper acetate called verdigris, which is used as a pigment, is made by exposing copper to acetic acid in the presence of air. The formula assigned to it is Cu(OH)2,2Cu(C2H 3 O 2 )2. Paris green is prepared by heating verdigris with a solution of arsenious acid containing acetic acid; the formula assigned to the light- green product formed is Cu(C2H3O2)2,Cu3(AsO3)2; it is used to exterminate potato bugs and insects. Malachite is a basic car- bonate, CuCO3,Cu(OH) 2 , found in nature, which is used as a green pigment. A compound of the same composition is formed when sodium carbonate is added to a solution of a copper salt; the reaction resembles that which takes place in the case of zinc and is COPPER, SILVER, AND GOLD 599 due to the same cause (665). Cuprous sulphide, Cu2S, is formed when copper is heated with sulphur, and cupric sulphide, CuS, when hydrogen sulphide is passed into a solution of a copper salt. 729. Tests for Copper Salts. Copper is precipitated in the second group in qualitative analysis since its sulphide, which is black, is precipitated by hydrogen sulphide in the presence of dilute acids. Copper sulphide dissolves in hot dilute nitric acid. Soluble cupric salts are blue in solution and give with ammonium hydroxide greenish precipitates, which dissolve in an excess of the reagent and produce deep blue solutions (595) . Cupric salts react with a solution of potassium ferrocyanide and give cupric ferrocyanide, which is a red-brown precipitate: 2CuS0 4 + K 4 Fe(CN) 6 = Cu 2 Fe(CN) 6 + 2K 2 S0 4 Copper compounds color the borax bead green in the oxidizing flame and red in the reducing flame. GOLD Gold is univalent in some of its compounds, which resemble the analogous derivatives of copper, and trivalent in others, in which it functions as a weak acid-forming element and resembles somewhat aluminium when the latter acts in this capacity. 730. Occurrence and Metallurgy of Gold. Gold occurs chiefly in the metallic condition either in the form of small particles mixed with sand, which has been washed into the beds of rivers, or dis- seminated through veins of quartz rock which have not been subjected to erosion. It occurs in this form in the Klondike, the Transvaal in South Africa, and California. It also occurs in small quantities in the sulphides of copper and lead, and is recovered from these metals in their purification (582). A mineral com- posed of the tellurides of gold and silver, (Au,Ag)Te2, is mined in Colorado; the metals obtained by heating it to drive off the tel- lurium are separated by treatment with strong sulphuric acid, which dissolves the silver. Since the acid will not attack the alloy unless it contains about three-fourths silver, this method of " parting " gold and silver is called " quartation." The silver sulphate formed is dissolved from the mixture with hot water, and the metal is precipitated by means of metallic copper. The gold 600 INORGANIC CHEMISTRY FOR COLLEGES left as a powder is melted in a cupel with potassium nitrate and borax, if base metals are present. 731. Native gold is extracted from the mass of rock with which it is mixed by amalgamation or cyaniding. The rock is first broken into small pieces and is then powdered in a stamp-mill, which consists of a series of iron mortars supplied with cylindrical pestles that are raised and lowered by machinery. The crushed ore, mixed with water, is fed into the mill continuously and escapes through a fine screen containing holes about 0.5 mm. in diameter. It next passes over inclined sheets of copper coated with mer- cury, from which the amalgam is scraped once a day. The latter is then pressed in chamois skin through which most of the mercury passes; the hard amalgam left is distilled in a retort. The residue is refined by heating with niter and borax, and " parted." In the cyanide process the finely pulverized ore is treated with a 0.25 per cent solution of sodium cyanide, which dissolves the gold in the presence of air: 4Au + SNaCN + O 2 + 2H 2 = 4NaAu(CN) 2 + 4NaOH The solution is next passed through boxes containing zinc shavings, or is agitated with zinc dust, and the precipitate formed treated with dilute sulphuric acid to dissolve the excess of zinc and the other base metals present. The product is cupelled and parted. Gold is sometimes obtained by chlorination. In this process the ore is treated with water containing chlorine and the metal is precipitated by ferrous sulphate from the chloride which is formed: 2AuCl 3 + 6FeS0 4 = 2Au + 2FeCl 3 + 2Fe 2 (S0 4 ) 3 or as a sulphide by hydrogen sulphide : 2AuCl 3 + 3H 2 S = Au 2 S 3 + 6HC1 The sulphur is removed from the compound by heating it. 732. Properties of Gold. Pure gold is bright yellow, but the color of native gold, which usually contains more or less silver, varies with the locality from which it is obtained. It is the most malleable and ductile of the metals, and can be made into sheets 0.0002 mm. thick; in this form it is used as gold-leaf in making signs and in lettering book covers. As gold is quite soft it is alloyed, COPPER, SILVER, AND GOLD 601 usually with copper, to make it more durable. The purity of gold is designated in carats, twenty-four being the number assigned to the pure metal. Jewelry is often made from 18-carat gold, which contains ^f of the metal, but on account of the softness of this alloy, 14-carat gold is usually preferred. The gold coins of Great Britain are made from 22-carat gold, the other metal in the alloy being copper. Similar alloys are used in the coins of the United States, France, and Germany, but they contain 90 per cent of gold, which is equivalent to 21.6 carats. In some countries coins are made from an alloy of gold and silver. The fineness of gold can be roughly determined by rubbing it gently with a fine- grained stone and testing with strong nitric acid the streak of metal formed; the result is compared with that produced with streaks made from gold of known composition. Gold is not affected by the substances in the air and does not decompose hydrogen sulphide. It does not dissolve in the com- mon acids, but is soluble in aqua regia as the result of the forma- tion of a soluble chloride. 733. Compounds of Gold. Auric chloride, AuCls, is a red crystalline compound formed by gently heating chlorauric acid, HAuCU, which is made by dissolving gold in aqua regia. A num- ber of double salts are derived from the acid; the yellow sodium salt, NaAuCU,2Il2O, is used in photography (719). Auric hy- droxide, Au(OH)3, is formed as a yellowish brown precipitate when a solution of an auric compound is treated with sodium hydroxide; it dissolves in a solution of sodium hydroxide and forms a salt of the composition NaAuO2, which is derived from metauric acid, Au(OH)3 B^O = HAuCb. Aurous chloride, AuCl, is formed by heating auric chloride at about 180; it is colorless and is converted in boiling-water into auric chloride and gold. Aurous iodide is precipitated when potas- sium iodide is added to a solution of auric chloride. Aurous oxide, Au2O, is a violet powder which reacts with hydrochloric acid to form auric chloride and gold. It will be seen that the aurous compounds resemble closely those of cuprous copper. Auric oxide, Au2Oa, is a brown powder. Hydrogen sulphide precipitates from a solution of auric chloride a brown mixture of aurous sulphide, Au2$, and auric sulphide, The sulphides dissolve in ammonium sulphide and form 602 INORGANIC CHEMISTRY FOR COLLEGES thioaurites and thioaurates. Potassium cyanide forms two classes of complex salts, the aurocyanides, like KAu(CN)2, and the auricyanides, of which KAu(CN)4 is an example. 734. Analytical Reactions of Gold. The behavior of soluble gold compounds with hydrogen sulphide puts the metal in the analytical group with arsenic, antimony, and tin; the sulphides of these metals are precipitated by hydrogen sulphide in the presence of strong acids, and dissolve in yellow ammonium sul- phide. A solution of gold chloride yields a characteristic precip- itate of the metal, called purple of Cassius, when it is treated with a dilute solution of stannous chloride. EXERCISES 1. How could you (a) show the presence of Ag in an aqueous solution of AgCl, (6) separate AgCl from Agl, and (c) prove that AgBr contained Br? 2. State three ways of distinguishing AgNO 3 from KNO 3 without the use of chemicals. 3. How could you separate pure Ag from a mixture of AgNO 3 and Pb(NO 3 ) 2 ? 4. Write equations for reactions by which Cu can be obtained from CuSCX by three different methods. 5. If a piece of silver upon which has been deposited a thin layer of Ag2>? is brought into contact with a piece of aluminium under a dilute solution of sodium hydroxide, what would you expect to happen? For what purpose could the reaction which takes place be used? 6. How could the coating that forms on copper in the air be removed? 7. How could you obtain in two different ways pure copper from a solution which contains CuSO 4 and A1 2 (SO 4 ) 3 ? 8. How could you distinguish by chemical means Cu from Cu 2 O? 9. Explain what you think would happen when a concentrated solution of hydrobromic acid is added to an aqueous solution of CuBr 2 . 10. How could you show the presence of Sb and Au in the sulphides pre- cipitated by H 2 S, which dissolve in (NH 4 ) 2 Sz? 11. When a sample of an alloy of silver and copper weighing 0.8572 gram was dissolved in nitric acid and the silver was precipitated as chloride, the latter weighed 0.7167 gram. Calculate the percentage of silver in the alloy. 12. Calculate what weight of (a) silver chloride and (6) 6H 2 can be obtained from 1 gram of silver obtained from a dime, CHAPTER XLI IRON, COBALT, AND NICKEL 735. Iron was used in the earliest times of which we have any historical records; furnaces for the metal are mentioned by Moses, and the art of making weapons from iron was known to the Egyp- tians and Hindus. The introduction in 1713 of coke in place of charcoal as the agent used to reduce iron ore, the utilization of a blast of air in 1823 to assist in obtaining the high temperature required in the furnace, and the process of making steel invented by Bessemer in 1856 are, perhaps, the three most important devel- opments in the iron industry which have given it such a com- manding position in modern life. 736. Occurrence and Metallurgy of Iron. Iron is found in the metallic condition in meteorites, which contain, in addition, from 3 to 8 per cent of nickel associated in certain cases with smaller amounts of cobalt and copper. The metal is widely distributed in nature in the form of silicates, oxides, and the sulphide, FeS2, and is the fourth in abundance in the earth's crust. The chief ores of the metal are hematite, Fe 2 03, limonite, 2Fe203,3H 2 O, magnetite, Fe 3 O4, and siderite, FeCOs. Iron occurs in plants and animals as a constituent of complex organic compounds; it is present in the hemoglobin of the blood, and is involved in this condition in the absorption of oxygen in the lungs. 737. Iron is obtained by reducing its ores with carbon. The ores contain, in addition to the oxides and carbonate of the metal, small amounts of combined sulphur, phosphorus, and man- ganese, and are mixed with more or less sand and clay. In order to convert the latter into compounds which are liquid at the tem- perature at which reduction is carried out and thus make it pos- sible to separate them from the metal produced, the ore is mixed with limestone; this furnishes the basic material to flux the silica 603 604 INORGANIC CHEMISTRY FOR COLLEGES and alumina by converting them into calcium silicate and calcium aluminate. The reduction of the ore is carried out in a blast furnace, which is charged with a mixture of the ore, calcium car- bonate, and coke. Blast furnaces (Fig. 40) are constructed of steel, lined with brick, and are from 80 to 100 feet high, and from 18 to 22 feet in diameter. The charge is carried by a mechanical conveyor to the top of the furnace and introduced into it through a hopper so constructed that when the operation is carried out there is little loss of gas from the furnace. A blast of highly FIG. 40. heated air enters the furnace through pipes, called the tuyeres, and rising through the charge converts the carbon near the bottom into carbon dioxide. The latter is subsequently reduced by the hot carbon above it to carbon monoxide, which passes up through the furnace and reduces the oxide of iron. 738. There are four zones within the furnace where different re- actions occur. (Fig. 41 .) In the upper zone, where the temperature is from 100 to 300, the moisture and most of the combined water in the ore are removed; in the next at 300 to 800, the oxides of iron are reduced by carbon monoxide, to a mixture of the metal and ferrous oxide. The reaction which takes place between an oxide IRON, COBALT, AND NICKEL 605 of iron and carbon monoxide, for example, FeaCU + 4CO = 3Fe + 4CO2, is a reversible one and exothermic, and, as a consequence, the equilibrium is displaced unfavorably for the reduction with rise in temperature. For this reason a large excess of carbon monoxide must be used and the part which is not oxidized by the iron escapes with the other gases from the top of the furnace. In Coke ^ i' me stone Ore Oxygens Nitrogen-* JOSCO ,3/eOi-COi \ Blowing HofKasf c ~ Stains Coke. Liquid Iron and Slqg_ _ ~ Cement rfil/zer. Mineral Woor. Wrouyfrffroo. FIG. 41. this zone the limestone and siderite, if it is present, are converted into oxides and carbon dioxide. In the third zone, at 800 to 1300, the reduction of the fer- rous oxide, FeO, is completed and the manganese, silicon, and phosphorus are converted by the hot carbon into the elementary condition and unite with the iron. In the zone of fusion, which comes next, the temperature is from 1300 to 1500. Here the 606 INORGANIC CHEMISTRY FOR COLLEGES iron melts and flowing over the carbon becomes saturated with it, and finally drops into the bottom of the furnace. In this zone the lime unites with the siliceous material and the calcium silicate and calcium aluminate produced melt and form the slag, which floats on the molten iron. 739. Blast furnaces are run continuously for months; the slag is tapped off through a hole at intervals of about two hours and the iron every four to six hours. The composition of a typical slag is 30 to 35 per cent SiO2, 10 to 15 per cent A^Oa, and 50 to 55 per cent CaO. Some slags are ground and used in cement-making, and others not adapted to this purpose are used as ballast on rail- roads, for road-building, and for other purposes. The molten iron is either converted directly into steel or run into molds and cast into " pigs." The gas which issues from the top of the furnace through a pipe called the " downcomer " contains over 20 per cent of carbon monoxide, and possesses a high heat-producing value. A part of it is used to heat the stoves through which the air-blast passes before being admitted to the furnace, and the rest is freed from dust and used in gas engines as a source of power. 740. Cast Iron. The metal obtained from the blast furnace is known as pig iron; it contains, in addition to the metal, from 3 to 4 per cent carbon, 1 to 3 per cent silicon, about 0.7 per cent each of manganese and phosphorus, and from 0.02 to 0.05 per cent sulphur. The properties which are characteristic of cast iron are due largely to the high percentage of carbon that it contains. In the liquid metal the carbon is in solution in the form of a carbide, Fe 3 C, which is called cementite. The properties of the solid iron are determined by the rate at which the metal cools. If this takes place slowly, as in a sand mold, some of the carbide decomposes into iron and carbon, which crystallizes in thin black scales as graphite, and gray cast iron is obtained. If the cooling takes place more rapidly, less carbon separates and white cast iron is formed. When the molten iron is cast in a mold made of metal the part next the mold is suddenly chilled and is found to consist of a solid solution of cementite in iron without any graphite. It is called chilled cast iron and is very hard and brittle. Phosphorus increases the fluidity of the molten metal and aids in casting it. IRON, COBALT, AND NICKEL 607 741. Wrought Iron. The properties of pure iron are very dif- ferent from those of cast iron; it softens before melting, which takes place at 1530, whereas cast iron melts more or less sharply at temperatures between 1150 and 1250, which depend on the amounts of the other elements present. The purest form of com- mercial iron is used in the manufacture of piano wire, which con- tains 99.7 per cent of the metal. Wrought iron, which is so called because it softens on heating and can be worked when hot, is made from cast iron by removing most of the other elements present in it. The pigs are melted on the hearth of a reverberatory furnace, which is covered with hematite or magnetite. The car- bon in the pig iron unites with the oxygen contained in the oxides and escapes as carbon monoxide. The phosphorus and silicon are also oxidized and pass into the slag as phosphates and silicates along with ferrous sulphide and any manganese in the pig iron. The mixture is stirred or " puddled," and as the change to pure iron progresses it becomes stiffer and stiffer, and finally is worked into balls weighing about 150 pounds, which consist of minute globules of iron mixed with slag. These are then rolled to squeeze out as much as possible of the slag and to make the particles of iron coalesce. This process leaves about 2 per cent of slag in the metal which contains less than 1 per cent of other impurities, of which about 0.2 is carbon. The presence of slag in the iron causes it to have a fibrous structure which is said to increase its resistance to breaking and to assist when the metal is welded. Wrought iron is now often replaced by soft steel, which resembles the former in composition closely but does not have the texture which results from the presence of slag. 742. Steel. Cast iron, as we have seen, is hard and brittle, whereas wrought iron is soft and can be welded; the first contains a relatively high, and the second a very small percentage of carbon: Steel stands between these two in properties and chemical com- position, and can be made of any desired degree of hardness by regulating its content of carbon. Soft steels that can be welded contain about 0.2 per cent of carbon, and those designed for struc- tural work from 0.2 to 0.6 per cent, the low carbon content per- mitting the rolling of the metal into the desired form. The hard steels required for tools contain from 0.9 to 1.5 per cent of carbon. The properties of iron are influenced by the elements other than 608 INORGANIC CHEMISTRY FOR COLLEGES - Flame carbon in pig iron, and in the preparation of steel from the latter most of the sulphur, phosphorus, and silicon are removed. In converting pig iron into steel, the metal is largely freed from the other elements present by oxidizing them in the presence of substances which unite with the oxides formed to produce a slag. The desired amount of carbon is then added along with other sub- stances which are used to remove the last traces of oxygen or oxides present, or to give the steel the special properties desired. A number of ways of accomplishing this are in use, the more important of which will be described briefly. 743. The Bessemer Process. In this process as used in the United States molten pig iron is poured into a converter (Fig. 42) which is constructed of steel, lined with sand and clay, and supported on trunnions so that it can be rotated to pour out the finished material. Cold air is blown through the molten metal and the elements present in the iron are converted into oxides; the silicon burns first and unites with the man- ganese oxide formed and some iron oxide to produce silicates, which sepa- rate as a slag; the carbon is next oxidized to carbon monoxide which burns at the mouth of the converter. Much heat is generated during the process and the temperature is kept down by adding solid scrap iron or by introducing steam with the blast. In about fifteen minutes, when carbon monoxide no longer burns at the mouth of the converter, an alloy of manganese, iron, and carbon is added in the amount necessary to produce a steel of the required com- position. Sulphur and phosphorus are not removed from iron in the Bessemer process, because at the temperature of the converter iron has a greater affinity for these elements than oxygen. To make steel low in sulphur and phosphorus it is necessary to use pig iron containing not more than 0.1 per cent of each of these elements, or to line the converter with basic substances like ignited dolo- mite, which will take up sulphur and phosphorus from iron. Basic linings are used in Europe, but in the United States steel is made FIG. 42. IRON, COBALT, AND NICKEL 609 from pig iron high in phosphorus by the open-hearth method which is described below. 744. The Open-hearth Process. In this process pig iron is heated on a hearth the lining of which is determined by the amount of the elements to be removed from the iron. If the metal is low in phosphorus and sulphur, the lining is made of sand or other siliceous material, and the chemical reactions that take place resemble those in the Bessemer converter; this is known as the acid open-hearth process. If the pig iron contains more of these elements than can be present in the steel, the lining is made of calcined magnesito or dolomite, which takes up the phosphorus and sulphur from the iron. The basic open-hearth process is commonly used in the United States. FIG. 43. The hearth is covered with a roof made of silica brick (Fig. 43) and is heated by gas which burns directly above it. The charge, which consists of pig iron and enough hematite or other oxide of iron to furnish the required amount of oxygen to convert the sulphur, phosphorus, and silicon into oxides, is put on the hearth and heated for about ten hours. In order to obtain as high a tem- perature as possible the gases are preheated before they are brought together over the hearth. This is accomplished by using what is called a regenerative furnace in connection with the hearth. The hot products of combustion are passed through two chambers containing " checker-work" made of fire-brick, until they become red hot. Gas is next passed through one of the chambers and air through the other and the two brought together over the hearth. 610 INORGANIC CHEMISTRY FOR COLLEGES The products of combustion are passed through a second pair of chambers filled with checker-work. By shifting at intervals the supply of air and gas from one pair to the other, it is possible to prevent the loss of the large amount of heat which under ordinary conditions passes off with the products of combustion. The production of steel in this way is slow and the composition of the metal can be tested from time to time by the analysis of a sample taken from the hearth. When the desired changes have taken place, the material is drawn off into ladles to which are added the material to give it the desired composition. Aluminium is often used to remove the dissolved and combined oxygen, and the oxide formed passes into the slag. Titanium is also used, because it unites with both the oxygen and the dissolved nitrogen. The steel made with these deoxidizers is less apt to break under strain, and is now commonly used in making rails. 745. Crucible Steel. In manufacturing certain kinds of steel the process is carried out in crucibles made of a mixture of graphite and fire-clay in equal amounts. The charge consists usually of a low-carbon steel or wrought iron and an alloy of iron and manga- nese, either f erromanganese or spiegel-eisen (788) . If other metals are desired in the steel, such as nickel, tungsten, chromium, vana- dium, and molybdenum, they are also added. The crucibles are covered and heated for about four hours. Steels prepared in this way are used in making tools, drills, dies, files, springs, shafts, etc. Large quantities of steel are now made in electric furnaces. 746. Tempering of Steel. The physical properties of steel are determined not only by the amount of carbon and other ele- ments it contains but by the heat treatment to which it has been subjected. Steel which has been heated to redness and allowed to cool slowly is comparatively soft; if it is plunged into water when hot it becomes very hard. Any degree of hardness between the two extremes can be obtained by heating the chilled metal at different temperatures for some time. This process is called tempering, and the temperatures used vary with the purpose to which the steel is to be put. 747. The changes which take place in tempering steel are complex and result from the fact that pure iron can exist in three distinct modifications, and that the solubility of cementite, the carbide of iron, FesC, which is present in steel, varies with the IRON, COBALT, AND NICKEL 611 temperature. Further, the changes that take place occur slowly and it is possible by more or less rapid cooling to pass the transition points at which one form of iron changes to another; in the super-cooled condition of the metal which results, a variety of forms may exist. Pure iron is called ferrite, and its three modifications are indi- cated by prefixing to the name the Greek letters alpha,, beta, and gamma. a-Ferrite is the form of iron which is stable below 760. When pure iron is heated a change occurs at this temperature which becomes evident by the heat which is absorbed without affecting the temperature; it is a true transition temperature where a-ferrite changes into /3-ferrite. A second transition occurs at 900 when 7-ferrite is formed. The three forms differ in properties; the first can be magnetized, the other two not; 7-ferrite dissolves up to 1.7 per cent of carbon in the form of cementite, FeaC, whereas /3-ferrite dissolves much less and a-ferrite, none at all. If steel is heated above 900 and cooled rapidly, there is not time for the cementite to separate, and the product obtained is a solid solution of the compound in 7-ferrite, which is very hard. If the iron contains more than 1.7 per cent carbon, the excess separates as graphite, and the product is known as pig-iron. When the steel cools slowly there is sufficient time for the cementite to crystallize out and for the 7-ferrite to change to a-ferrite. Under these conditions the product is an intimate mixture of iron and cementite, and is comparatively soft if it contains less than 0.8 per cent of carbon. When steel is heated to temper it, its constituents change slowly. When the transformations have taken place to such a degree that the metal after cooling has the desired hardness, it is more or less rapidly chilled by introducing it into liquids which take up the heat; if the cooling is to be very rapid, it is plunged into water; if less rapid, into oils or molten salts. 748. Alloy Steels. By alloying steel with certain metals, products are obtained that possess properties which adapt them to special uses. The effect of the added metal is, in general, to lower the transition temperatures at which the different forms of iron pass one into the other; as a result, 7-ferrite and its solid solution of cementite, which is very hard, exist at lower temperatures than in ordinary steel. For example, manganese steel which contains 612 INORGANIC CHEMISTRY FOR COLLEGES from 12 to 14 per cent of the metal and about 1.5 per cent carbon, can be cast and, after rapid cooling, is very hard. It is used in making rock-crushing machinery, steam shovels, burglar proof safes, and for other purposes which require a very hard steel. The addition of a very small amount of tungsten or molybdenum increases the hardness of the metal. Nickel steel is used when exceptional toughness, strength, or hardness is desired and when a metal is required to resist abrasion or corrosion. The proportion of nickel in the steel varies from 3.5 to 42 per cent. Chromium steels are also hard and tough; they are used for files, safes, plows, etc. Very small amounts of vana- dium produce about the same effect on steel as that produced by nickel. Steels of this type contain approximately 0.15 per cent vanadium, 0.25-0.5 per cent carbon, 0.5-1 per cent manganese and 1 per cent chromium or nickel; they have a high elasticity, are hard and tough, weld readily, and resist shock and vibratory stresses. For this reason they are used in making automobile parts. High-speed steels do not lose their hardness and toughness at red heat, and are used in making tools for cutting metals at such a speed that the friction produced heats them to a high temperature. They are made from steel to which 15 to 20 per cent of tungsten, 3 to 5 per cent of chromium, and a small quantity of vanadium or cobalt have been added. An alloy of iron and silicon containing 14 to 15 per cent of the latter, which is sold under the trade name "duriron," is used in making castings for chemical apparatus that resists to a high degree corrosion by acids. 749. The Properties of Iron. The more important physical and chemical properties of iron have already been given. When iron burns in oxygen or is heated to a high temperature in steam, it is converted into an oxide of the composition FeaCU. The reaction with steam is utilized in covering iron with a coating which protects it against rust (552). Passive iron and the action of acids on the pure metal have been discussed (548, 547). Iron forms two well-defined series of compounds in which it acts as a base-forming element; the ferrous compounds, in which the metal has the valence 2, resemble the compounds of mag- nesium and are hydrolyzed to but a slight degree; the ferric com- pounds contain trivalent iron and closely resemble the analogous derivatives of aluminium. Iron acts as an acid-forming element IRON, COBALT, AND NICKEL 613 when it has the valence 6, and forms salts, called ferrates, which resemble in composition the sulphates. 750. Ferrous Compounds. When iron dissolves in dilute acids it is converted into ferrous salts. Ferrous chloride, FeCl2,4H2O, forms pale blue crystals, which turn green in the air as the result of oxidation which converts them into a basic ferric chloride. In solution the salt is also oxidized by air; it can be kept in the ferrous condition by having metallic iron and a small amount of hydro- chloric acid present. In order to obtain ferrous hydroxide in pure condition the solution from which it is precipitated by an alkali must be free from oxygen and ferric salts. This is best accom- plished by treating with a boiling solution of sodium hydroxide, a boiling solution of ferrous sulphate which contains a little sul- phuric acid and is in contact with metallic iron. Under these conditions the hydroxide is formed as a white precipitate. It changes rapidly in the presence of air to a green color, and finally is converted into ferric hydroxide, Fe(OH)3, which is brown. Ferrous carbonate, FeCOs, is obtained as a colorless precipitate when a soluble carbonate is added to a ferrous salt in the absence of oxygen. It is soon converted in the presence of air into ferric hydroxide as the result of oxidation to the ferric salt, which then hydrolyzes. Ferrous carbonate, like calcium carbonate, is soluble in water containing carbonic acid as the result of the formation of an acid salt, Fe(HCO3)2- It is present in this form in hard waters containing iron, which deposit ferric hydroxide as the result of the loss of carbon dioxide and subsequent oxidation when the water comes into contact with the air. Ferrous sulphate, FeSO4,7H2O, which is also called green vitriol or copperas, is obtained by evaporating a solution prepared by dissolving iron in dilute sulphuric acid. The salt is efflorescent and becomes brown in the air as the result of oxidation to a basic ferric sulphate, Fe(OH)SO4. The double salts containing ferrous iron are much more stable in the air than the simpler compounds. Ferrous-ammonium sulphate, (NBU^SO^FeSO^GH^O, crystal- lizes well from water and can be kept in a pure condition ; it is used in standardizing solutions for use in quantitative analysis. Ferrous sulphate is used as a mordant in dyeing, as a disinfectant, in the purification of water supplies, in the manufacture of ink and Prussian blue, and in metallurgy for precipitating gold. 614 INORGANIC CHEMISTRY FOR COLLEGES Ferrous sulphide, FeS, which is black, is made by heating iron and sulphur, and is used in making hydrogen sulphide. It is formed by adding ammonium sulphide to a solution of a ferrous salt. 751. Ferrous cyanide, Fe(CN) 2 , is precipitated when potassium cyanide is added to a ferrous salt. It dissolves in an excess of the reagent to form a stable complex cyanideof theformulaK4Fe(CN)6,- 3H 2 O, which is called potassium ferrocyanide or, in trade, yellow prussiate of potash. It is manufactured by heating scrap iron with potassium carbonate and animal refuse material which con- tains nitrogen and sulphur, such as hair, hoofs, blood, and leather scrap. Potassium ferrocyanide yields K + and Fe(CN)e~" ions when dissolved in water. The fact that it is not poisonous indicates that it does not yield any potassium cyanide. Its solu- tions give none of the reactions characteristic of ferrous salts. Potassium ferrocyanide precipitates insoluble ferrocyanides from solutions of salts of most metals. Copper ferrocyanide, Cu2Fe(CN) 6, obtained in this way is brown and serves to aid in the identifica- tion of copper salts. Ferrous ferrocyanide is white, but soon turns blue as the result of oxidation. Ferric ferrocyanide, Fe4[Fe(CN)e]3, is a deep-blue precipitate which is called Prussian blue and is used as a pigment and in making laundry blueing. Potassium ferrocyanide is used in dyeing and calico printing, and in making potassium cyanide and potassium ferricyanide (753). 752. Ferric Compounds. Ferric chloride in the anhydrous condition is prepared by heating iron in a stream of chlorine; it sublimes without decomposition and forms green scales. The salt obtained by treating ferrous chloride dissolved in water with chlorine and evaporating the solution, is a deliquescent hydrate of the composition FeCl3,6H2O. The yellow color of a solution of ferric chloride is largely due to the colloidal ferric hydroxide formed as the result of hydrolysis. The color can be intensified by heating the solution or by adding to it a small amount of sodium hydroxide to neutralize the free acid present; in both cases the hydrolysis is increased. Ferric salts are converted by reducing agents into ferrous salts. Stannous chloride or metallic zinc are commonly used for this purpose : 2FeCl 3 + SnCl 2 = 2FeCl 2 + SnCU 2FeCl 3 + Zn = 2FeCl 2 + ZnCl 2 IRON, COBALT, AND NICKEL 615 These reactions are used in volumetric analysis in the quantitative determination of iron. The metal is reduced to the ferrous con- dition, and the volume of a standardized solution of an oxidizing agent required to convert it into a ferric salt is determined. Ferric hydroxide is formed as a brown precipitate when a base is added to a solution of a ferric salt. It is not soluble in an excess of alkali. It is readily obtained in the colloidal condition by suspending in a stream of water a solution of ferric chloride con- tained in a bag made of parchment, which is an animal mem- brane. The molecules of hydrochloric acid which result from the partial hydrolysis of the chloride are of such size that they can pass through the minute pores in the parchment, but the larger mole- cules of the ferric hydroxide cannot. As the acid diffuses through the membrane the equilibrium between it and the hydroxide is destroyed FeCl 3 + 3H 2 ;= Fe(OH) 3 + 3HC1 and finally all of the salt is hydrolyzed. Prepared in this way ferric hydroxide remains in colloidal solution that is, it is apparently dissolved in the water, but it shows none of the properties of an electrolyte ; it has no effect on the freezing-point and does not conduct an electric current. The solution, known as dialyzed iron, has a red color and is used as an antidote in arsenical poisoning. Ferric hydroxide is precipitated when a carbonate is added to a solution of a ferric salt, the reaction being analogous to that which occurs in the case of aluminium salts. When solutions containing the salts of bivalent -and trivalent metals are shaken with solid barium carbonate, the hydroxides of the trivalent metals only are precipitated. The reaction serves, therefore, as a means of sep- arating these two kinds of metals and is used in qualitative analysis. Ferric oxide, Fe20s, is extensively used as a red pigment. It is prepared by treating with sodium carbonate the liquors obtained in cleaning the surface of iron, and igniting the precipitate formed. In this form it is sold under the name rouge. If sulphuric acid has been used for " pickling," the solution is first neutralized with lime; the precipitate, which contains ferric hydroxide mixed with calcium sulphate, yields on ignition a light red pigment called Venetian red. The oxide of iron of the formula FeaCU, which occurs in nature, is called lodestone or magnetic oxide on account of the fact that it 616 INORGANIC CHEMISTRY FOR COLLEGES will attract iron. It is formed by heating iron in oxygen or super- heated steam, and is a compound of ferrous and ferric oxides. Ferric sulphate is made by oxidizing with nitric acid a solution of ferrous sulphate containing sulphuric acid. It does not crystallize well and is very soluble in water. Like the sulphates of the other trivalent metals it forms alums. 753. The addition of a cyanide to a ferric salt causes the pre- cipitation of ferric cyanide, which dissolves in an excess of the reagent and forms potassium ferricyanide, K 3 Fe(CN)e [Fe(CN)3, 3KCN], which, on account of its color, is called red prussiate of potash. It is manufactured by treating a solution of potassium ferrocyanide with chlorine: 2K 4 Fe(CN) 6 + C1 2 = 2K 3 Fe(CN) 6 + 2KC1 When a solution of potassium ferricyanide is added to one of a ferrous salt, a precipitate, called Turnbull's blue, is formed: 3FeSO 4 + 2K 3 Fe(CN) 6 = Fe 3 [Fe(CN) 6 ] 2 + 3K 2 SO 4 With ferric salts the ferricyanide gives no precipitate, but a brown solution is formed. The use of the ferricyanide in photography has been mentioned (719) . 754. Iron as an Acid-forming Element. Ferric hydroxide does not dissolve in solutions of bases, but when it is fused with caustic alkalies it reacts with the base and salts, called ferrites, are pro- duced. These salts, of which sodium ferrite, NaFe02, is an exam- ple are derived from the compound formed from ferric hydroxide by the loss of water, Fe(OH)3 H 2 = HFeO 2 ; they are com- pletely hydrolyzed by water, which produces the base and ferric hydroxide: NaFeO 2 + 2H 2 O = NaOH + Fe(OH) 3 . This reac- tion has been utilized in a commercial process (Loewig's) for the manufacture of caustic soda and yields a very pure product. Iron oxide is heated at a high temperature with sodium carbonate and the product, after being washed with cold water, is decom- posed by hot water. The iron hydroxide is recovered and used again. If a solution of potassium hydroxide in which precipitated ferric hydroxide is suspended, is treated with chlorine, the iron is oxidized and converted into potassium ferrate, K 2 FeO 4 , which is obtained in the form of red crystals on evaporating the solution. IRON, COBALT, AND NICKEL 617 The salt is hydrolyzed by water in the absence of an excess of alkali, and the ferric acid formed decomposes into ferric hydroxide and oxygen. In the ferrates, which resemble the sulphates in composition, iron has the valence 6 as has sulphur in sulphates. The metal can be oxidized to this valence only in the presence of strong alkalies with which the acid formed can unite to form a salt. 755. The Corrosion of Iron. The changes that take place when iron rusts have been discussed briefly (550), but their fuller consideration was delayed until the properties of the oxides and hydroxides of the metal and the hydrolysis of its salts had been studied, because all these are involved in the process. On account of the industrial importance of the corrosion of iron, the changes involved in it have been carefully studied in order to arrive at an explanation of the process, upon which methods to prevent cor- rosion can be based. A number of views as to the way in which iron rusts have been advanced, but the one based upon electro- chemical considerations is now generally accepted. According to an older view carbon dioxide was necessary in ordinary rusting; it united with the water present to form carbonic acid, which reacted with iron and formed ferrous carbonate and hydrogen. The salt was next hydrolyzed by water, and the ferrous hydroxide produced was oxidized by the air to ferric hydroxide; the carbon dioxide liberated served to react with more iron, and the process continued. In the more recent explanation, the presence of carbon dioxide is not considered necessary, although it increases the rate of cor- rosion in the manner just indicated. In the presence of pure water iron passes into solution to a slight degree as ferrous hydroxide, and hydrogen is liberated. In the presence of the air the dis- solved hydroxide is oxidized to ferric hydroxide, which separates from solution on account of its very small solubility. Pure iron rusts very slowly, but commercial iron and steel, which contain a number of other substances, corrode rapidly. According to the electrochemical explanation of the process, the other substances present in the iron form with it an electric couple, of which the metal is the more positive constituent. Local action is set up in the way already described in the case of metallic couples (567), and the iron is rapidly converted into ferrous hydroxide, which is oxidized to the hydrated ferric oxide. Iron does not rust evenly, 618 INORGANIC CHEMISTRY FOR COLLEGES but at certain places on its surface, and pitting occurs where there is a segregation of the impurities in the metal. The presence of " mill scale/' which is the oxide formed on the surface of the hot metal when it is worked, markedly increases the rate at which iron rusts. The scale is brittle and when it cracks and leaves a surface of iron exposed, a couple is set up and the metal is cor- roded. Rust itself forms a couple with pure iron and when it is once formed the corrosion proceeds more rapidly. It has been shown that the rate of rusting during the second year is about twice as fast as during the first year. Unused rail- road rails rust more rapidly than those from which the rust is removed by the jarring produced by the passage of cars over them. When iron is strained in any way it assumes a different potential from that of the unstrained metal, and, as a result, corrosion is greater in the neighborhood of punched holes than around holes drilled in the metal. Corrosion takes place where the metal has been scratched by a file or struck with a heavy tool. Since oxygen is involved in the corrosion of iron, the metal rusts very slowly when it is immersed in deep water, which con- tains but a little of the gas. Rain water, which is saturated with oxygen and carbon dioxide, affects iron very rapidly. 756. Tests for Iron Salts. Solutions of ferrous salts give a deep-blue precipitate with potassium ferricyanide (753) and a black precipitate of ferrous sulphide, which is soluble in acids, on the addition of a solution of ammonium sulphide. Ferric salts give a deep-blue precipitate with potassium ferrocyanide (751) and an intense red coloration with ammonium thiocyanate, NH4SCN. When sodium acetate is added to a solution of a ferric salt, a red solution of ferric acetate is formed, which deposits, when heated, an insoluble basic acetate that results from the hydrolysis of the salt. The reaction is utilized in the quantitative separation of iron from manganese in the analysis of steel. The acetates of the trivalent metal are converted into insoluble basic acetates when heated in water solution, whereas those of the bivalent metals, which are stronger base-forming elements, are not decom- posed in this way. Iron salts color the borax bead green in the reducing flame (ferrous borate) and yellow in the oxidizing flame (ferric borate). IRON, COBALT, AND NICKEL 619 COBALT 757. Cobalt occurs as smaltite, CoAs2, and cobaltite, CoAsS, and is associated with nickel in its ores. The metal has been used recently in high-speed tool steels (748), and it has been suggested as a substitute for nickel in plating other metals. Cobalt forms two series of compounds in which it has the valence 2 and 3, respectively. The cobaltous compounds are but slightly hydro- lyzed and resemble those derived from bivalent iron. Cobaltous chloride, CoCl 2 ,6H 2 0, crystallizes from water in red prisms, which become blue when dehydrated. Solutions of caustic alkalies precipitate from solutions of cobaltous compounds blue basic salts that are converted by boiling into cobaltous hydroxide, Co(OH)2, which is pink, when first formed, but changes to brown as the result of oxidation to cobaltic hydroxide, Co (OH) 3. Cobaltous hydrox- ide dissolves in a solution of ammonia to form the compound Co(NH3)4(OH)2, which is rapidly oxidized to the cobaltic com- pound, Co(NHs)6(OH)3, which is blue. Cobaltous sulphate has the composition CoSO4,7H2O, and cobaltous nitrate, Co(NO3)2,- 6H2O; both salts are red. The sulphide, CoS, is black and is precipitated by ammonium sulphide. 758. When a solution of sodium hypochlorite is added to a solution of a cobaltous salt, cobaltic hydroxide is formed as a black precipitate. The hydroxide dissolves in cold hydrochloric acid with the formation of cobaltic chloride, which decomposes on warming into cobaltous chloride and chlorine. The cobaltic salts are all highly hydrolyzed in water, but are more stable in the form of double salts. Of these the complex cyanides are the most important. Potassium cobalticyanide, K3Co(CN)6, is prepared by the action of chlorine on potassium cobaltocyanide, K4Co(CN)e. The compounds resemble closely the analogous salts containing iron. A potash-cobalt glass is used under the name smalt as a pig- ment and in making blue glazes for china. It is prepared by fusing together sand, potassium carbonate, and cobaltic oxide, Co2Os. Cobalt blue is an excellent pigment made by igniting a mixture of alumina, and basic cobalt phosphate previously pre- pared by the action of sodium phosphate on a solution of cobalt nitrate. 620 INORGANIC CHEMISTRY FOR COLLEGES 759. Tests for Cobalt Salts. When cobaltous salts are treated in solution with acetic acid and potassium nitrite, they are first oxidized to the cobaltic condition by the nitrous acid set free, and are then converted into a complex nitrite of the formula KsCo(NO2)6, which is formed as a yellow precipitate. The reaction is characteristic of cobalt and is used as a test for the element in qualitative analysis. Cobalt colors the borax bead blue in both the oxidizing and reducing flame. NICKEL 760. Nickel occurs in the metallic condition in meteorites. The chief source of the world's supply of the metal is the Sudbury district in Ontario, where an iron sulphide occurs which contains about 2 per cent each of nickel and copper. A hydrated silicate of magnesium, nickel, and iron is used as an ore in New Caledonia. The ores are smelted and the nickel and copper separated. In the Monde process the reduction is carried out in a Bessemer converter, and the nickel is then converted into nickel carbonyl, Ni(CO)4, by heating the metal at about 50 in a stream of carbon mon- oxide. The product, which is a gas, is passed through cylinders heated to 200 in which nickel is deposited as the result of the decomposition of the carbonyl at this temperature. Nickel is separated electrolytically from copper, or the alloy which is the result of the reduction of the ore is used as such under the name monel metal, which is stronger than ordinary steel and resists the action of acids. Nickel is used in alloys (542), for plating (583), and in the production of certain steels; about 60 per cent of the world's production is used for the latter purpose (748). Invar, a steel which contains 30 per cent nickel, has a heat-coefficient of expan- sion of practically zero and is used in making pendulums, etc. Nickel coins contain 1 part of the metal alloyed with 3 parts copper. 761. Nickel forms two oxides, NiO and Ni20a. Nickelous hydroxide, Ni(OH)2, is converted by acids into nickelous salts, which are stable in the air and are not oxidized to nickelic com- pounds. Nickelic hydroxide, Ni(OH)s, has even less basic pro- perties than cobaltic hydroxide; it is formed as a black precipitate when a hypochlorite is added to a solution of a nickel salt. IRON, COBALT, AND NICKEL 621 The chloride, NiCl 2 ,6H 2 0, and sulphate of nickel, NiS0 4 ,7H 2 O, are green salts which crystallize from water. The double sulphate of the formula (NH 4 ) 2 Ni(SO 4 ) 2 ,6H 2 O is used in nickel plating (748). Nickelous hydroxide, which is green, is obtained by precipitation; it dissolves in ammonia and forms a colorless solution of a com- pound of the formula Ni(NH 3 ) 4 (OH) 2 . 762. Test for Nickel Salts. When potassium cyanide is added to a solution of a nickel salt the green nickelous cyanide, Ni(CN) 2 , first precipitated dissolves in an excess and forms a complex cyanide of the formula K 2 Ni(CN) 4 ,H 2 O, which does not resemble in composition and properties the double cyanides containing iron and cobalt in the bivalent condition. It is less stable and yields black nickelic hydroxide when treated with hypochlorites. The reaction serves to separate nickel from cobalt since potas- sium cobalticyanide is not affected by hypochlorites. Nickel differs from cobalt in that it does not form an insoluble double nitrite. EXERCISES 1. How could you distinguish from one another the following ores of iron: (a) hematite, (6) limonite, (c) magnetite, and (d) siderite? 2. What would happen if oxygen -were passed over metallic iron con- tained in a silica tube that was heated red hot? 3. A large excess of CO is required in reducing the oxides of iron in a blast furnace. What conclusion can you draw in regard to the reaction FeO -f- CO = Fe + CO 2 ? Is the conclusion in accord with the facts? 4. What chemical compounds would you expect to be present in a blast- furnace slag? 5. Name some articles made of (a) cast iron, (6) wrought iron, (c) high carbon steel, and state a reason in each case. 6. Write equations for the reactions which occur in (a) the acid Bessemer process, (6) the basic open hearth process. 7. (a) Why cannot FeCl 3 be made by heating FeCl 3 ,6H 2 O? (6) What would be obtained if the latter salt were heated to a high temperature? 8. Write an equation for the reaction which takes place when an aqueous solution of FeSO 4 containing H 2 SO 4 is heated with HNO 3 . 9. How could you test (a) a ferric salt for the presence of a ferrous salt, and (6) a ferrous salt for the presence of a ferric salt? Write equations for the reactions used. 10. Could Fe 2 O3 which contains A1 2 O3 be used in Loewig's process for making NaOH? Give a reason for your answer. 11. Write an equation for the reaction which occurs when K 2 FeO< is dissolved in water. 622 INORGANIC CHEMISTRY FOR COLLEGES 12. Is it a good practice to set iron fence posts in stone by placing the post in a hole in the stone and filling the cavity with lead? What improve- ment could you suggest? 13. What would be the result if an iron pipe laid in the ground remained in contact for a long time with a wire carrying an electric current? Explain. 14. (a) Out of what form of iron should anchors be made? Give reasons. If an iron anchor is used to hold a mooring in place and remains under water a long time what would you expect if (6) the water is shallow and the bottom is rocky, (c) the water is shallow and the bottom is covered with growing plants, (d) the water is deep and the bottom covered with soft mud? 15. Write equations for the reactions which take place when a solution of ferric chloride is treated with sodium acetate and then heated. Assume that the basic acetate formed has the formula Fe(OH) 2 (C.,H 3 O 2 ). 16. Show by a consideration of the properties of the compounds of Ni and Co which of the elements is the more closely related to Fe. Are the atomic weights of the elements in accord with your conclusion? CHAPTER XLII THE PLATINUM METALS 763. The metals in the eighth group in the periodic classification other than iron, cobalt, and nickel are classed together under the name platinum metals, because they are found associated with platinum and resemble the latter in physical properties. They occur as alloys which frequently contain more or less gold, and are obtained as nuggets and small particles by washing the alluvial sands deposited by certain rivers. The world's supply is obtained almost exclusively from the Ural Mountains, although small amounts of platinum are separated from the gold found in Australia, Brazil, and California, and from the copper-nickel-iron deposits of Sudbury, Ontario. 764. Platinum. In separating platinum, which is the most important metal of the group, the naturally occurring alloys are treated with aqua regia, and to the solution of the chloride formed is then added ammonium chloride, which precipitates a double chloride of the formula (NH^PtCle- This salt on being strongly heated decomposes and leaves a residue of platinum. Platinum is a soft, very heavy metal which has the specific gravity 21.4 and melts at 1755. Its resemblance to silver in appearance led the Spaniards, who discovered it in South America, to give it the name platinum, which is the diminutive of the Spanish word for silver, plata. The fact that the metal is com- paratively inactive chemically has adapted it to many uses. It has recently been used extensively in making jewelry, and it is claimed that it accentuates the beauty of diamonds when they are set in it, but its popularity can, no doubt, be traced, in part, to the fact that at present it is much more costly than gold. The most important uses of the metal are in chemical industries and in laboratories. In the manufacture of sulphuric acid it finds extensive application as a contact agent in the preparation of sulphur trioxide, and in the construction of evaporating pans 623 624 INORGANIC CHEMISTRY FOR COLLEGES used in making the concentrated acid. It is the most important catalyst used in the. oxidation of ammonia to nitric oxide, which is converted into nitric acid. The metal is used in the laboratory in the form of resistance wire, electrodes, crucibles, dishes, etc. It was formerly used in large quantities in the construction of electric-light bulbs and other electrical appliances in which it is necessary to pass a metallic conductor through glass. The variation of the coefficient of expansion of platinum with temper- ature is about the same as that of glass, and, as a consequence, the latter does not crack where the metal has been sealed into it. Platinum has been largely replaced for this purpose by wire made of nickel-steel coated with copper. 765. Platinum forms compounds in which it shows the valencies 2 and 4. When it has the higher valence it has acid-forming properties. For this reason the metal is attacked when in contact with fused alkalies, especially if oxidizing agents are present. It is not affected by fused alkaline carbonates and, consequently, crucibles of paltinum are used in the laboratory in preparing sub- stances for analysis which are decomposed by molten carbonates. Red-hot platinum allows hydrogen to pass through it freely, and unites with arsenic, lead, phosphorus, etc. For this reason sub- stances that contain elements which are reducible at red heat by hydrogen should not be ignited in a platinum crucible. Platinum also unites slowly with carbon at high temperatures and becomes so brittle that it is apt to crack. 766. When platinum is dissolved in aqua regia it is converted into chloroplatinic acid, H^PtClejOH^O, which forms reddish-- brown deliquescent crystals. The acid yields an ammonium salt, (NH^PtCle, and a potassium salt, K^PtCle; the latter, which is difficultly soluble in water, is used in the quantitative determi- nation of potassium. When chloroplatinic acid is heated at about 240 it is converted into platinous chloride, PtCl2, which is a green compound, insoluble in water, that dissolves in hydro- chloric acid and forms chloroplatinous acid; the potassium salt, K^PtCU is used in photography (719) . The double platinocyanides of potassium and barium, K 2 Pt(CN) 4 ,3H 2 and BaPt(CN) 4 ,4H 2 O, are employed in making screens for X-ray work; paper coated with either of these salts glows brightly when exposed to the rays on account of the fact that the compounds fluoresce, that is, they THE PLATINUM METALS 625 convert the invisible radiation of short wave-lengths into those of longer wave-lengths that produce the sensation of light. 767. Bases precipitate as a black powder from solutions of platinous chloride, platinous hydroxide, Pt(OH) 2 , which dissolves in acids but not in bases. Platinic hydroxide, Pt(OH)4, which is yellow, is formed in the same way from platinic compounds; it is soluble in both acids and bases, and with the latter yields plati- nates. Both sulphides of the metal, PtS and PtS 2 , dissolve in ammonium polysulphide and form ammonium thioplatinates. 768. The Other Metals of the Platinum Group. A large number of the compounds of the metals associated with platinum have been described, but cannot be considered here. Ruthenium and osmium form oxides of the formulas RuCU and OsC>4, and show, therefore, the valence 8, in accordance with their position in the periodic classification. The elements have acid-forming prop- erties and yield salts of the composition K 2 RuO4 and K 2 OsO4. The former, potassium ruthenate, is converted by large amounts of water into an oxide of the metal and potassium perruthenate, KRuO4, in which ruthenium has the valence 7. The tetroxide of osmium, OsCU, is commonly called " osmic acid/' although it does not form salts; it is used in histology for staining and hardening tissues. Osmium melts at about 2700 and was formerly used in making filaments for electric lamps but has been replaced for this purpose by tungsten. Wires made of an alloy of platinum and rhodium are used in one form of thermocouples. Rhodium is also used in making tips for gold pens. An alloy of iridium and osmium is also used for the latter purpose. Iridium is used in wires for thermocouples and is usually present in commercial platinum. The latter when pure is a relatively soft metal and a small per- centage of iridium makes it harder and, therefore, more durable. Palladium is characterized by its ability to absorb large quan- tities of hydrogen and is, therefore, a valuable catalytic agent in the union of the gas with other substances which react with hydro- gen. One volume of the metal in the form of foil will absorb 500 volumes of the gas, and when precipitated as a fine powder, about 800 volumes. The most characteristic compounds of the elements of this group are the double chlorides which they form with the chlorides of the alkali metals. CHAPTER XLIII CHROMIUM, MOLYBDENUM, TUNGSTEN, AND URANIUM 769. The consideration of the members of the first family of the sixth group in the periodic classification of the elements has been delayed to this point on account of the fact that the most important member of the family, chromium, forms compounds in which it plays the parts of a bivalent and a trivalent metal and others in which it acts as an acid-forming element. The chemistry of the compounds of chromium can now be studied in the light of the facts learned in regard to the behavior of bivalent and trivalent metals. The chromates, of which potassium chromate, E^CrCU, is an example, resemble closely in properties the sulphates. The similarity in their physical structure is evident from the fact that potassium chromate is isomorphous with potassium 'sulphate, and sodium chromate, Na2CrO4,10H2O, with Glauber's salt, which is the hydrate of sodium sulphate of analogous composition. It will be recalled that sulphuric acid and sulphur trioxide are oxidizing agents on account of the fact that the valence of sulphur can change from 6 to 4 or even less. In the case of chromic acid and its anhydride, CrOs, the reduction in valence takes place so readily that these compounds find many applications in chemistry as active oxidizing agents. When acting in this way chromium changes from the valence 6, in which it is an acid- forming element, to the valence 3, in which it is metallic in chemical properties and resembles aluminium. The other elements of the family form trioxides, which are acid anhydrides, and their most important compounds are salts which resemble the sulphates in composition. CHROMIUM 770. The most important ore of chromium is chromite, or chrome-iron ore, FeO,Cr2O3, which resembles magnetite, 626 CHROMIUM, MOLYBDENUM, TUNGSTEN, AND URANIUM 627 FeO,Fe 2 O3, in composition; it is considered to be a ferrous salt of an acid derived from chromic hydroxide by the loss of water Cr(OH)s H2O = HCrO2 and its chemical name is ferrous chromite. The element also occurs as crocoisite, which is lead chromate, PbCrC>4, and in traces in emeralds and other precious stones, to which it gives their characteristic colors. Chromium may be obtained by reducing the trioxide, Cr 2 O3, with aluminium, but the large amount of the metal used in making steel (748) is obtained by the reduction of chromite by carbon in the electric furnace; the alloy of chromium and iron obtained in this way is used directly. 771. Properties of Chromium. The element is crystalline, white, very hard, melts at 1520, boils at 2200, and has the specific gravity 6.9. It is not affected by the air, but burns in oxygen and forms the trioxide, Cr 2 O 3 . It dissolves slowly in hydrochloric acid and chromous chloride, CrCl 2 , is formed; its behavior with nitric acid resembles that of iron, and like the latter it is rendered passive by the concentrated acid (548) . The chief use of metallic chromium is in the preparation of alloys (748). 772. Chromous Compounds. Chromous chloride can be made as stated above, or by heating chromic chloride in hydrogen. It is also formed by the action of metallic zinc on a solution of chromic chloride: 2CrCl 3 + Zn = 2CrCl 2 + ZnCl 2 . The anhydrous salt is white and its solution in water is blue. It is very unstable and in the presence of hydrochloric acid is rapidly oxidized to chromic chloride by the oxygen of the air: 4CrCl 2 + 4HC1 + O 2 = 4CrCl 3 -f 2H 2 O. Chromous hydroxide, formed by the action of bases on chromous salts, is a yellow precipitate, which is such a powerful reducing agent^ that it reacts slowly with water and liberates hydrogen: 2Cr(OH) 2 + 2H 2 O = 2Cr(OH) 3 + H 2 . Chromous sul- phate, CrSO4,7H 2 O, resembles closely in chemical and physical properties ferrous sulphate, FeSO4,7H 2 0, but it is a much more active reducing agent. 773. Chromic Compounds. The large number of salts derived from trivalent chromium resemble closely those of aluminium and ferric iron. Chromic hydroxide, Cr(OH)s, is formed by the action of bases on chromic salts; it is pale blue in color and reacts with acids to form chromic salts, and with a large excess of a caustic alkali to form a chromite; potassium chromite, KCrO 2 , formed in 628 INORGANIC CHEMISTRY FOR COLLEGES this way is derived from an acid produced as the result of the loss of water from the hydroxide (754). The chromites are highly hydro- lyzed, and when heated with water are converted into chromium hydroxide, which is precipitated: KCrO 2 + 2H 2 = Cr(OH) 3 + KOH. Chromic hydroxide is formed when chromium salts are used as mordants (688) and in tanning (688), and is the compound which brings about the desired results in these important processes. Chromic chloride) CrCl3,6H 2 O, is obtained in the form of bluish- gray crystals by the evaporation at room temperature of a solution of chromic hydroxide in hydrochloric acid. When dissolved in water the salt gives a violet solution, which changes to green when heated to boiling. 774. From this solution can be isolated a second form of chromic chloride, which is green but has the same composition as that of the blue salt. Silver nitrate precipitates all the chlorine from the blue salt, but only one-third of that contained in the green chloride. This difference in behavior is explained on the hypoth- esis that the blue chloride gives the ions [Cr(H 2 O)e] + + " and 3C1~ and the green chloride the ions [CrCl 2 (H 2 O)4] + and Cl~. 775. Chromic oxide, C^Os, is used as a pigment under the name chrome green. It is prepared by igniting at red heat pre- cipitated chromic hydroxide. Chrome alum, K 2 SO4,Cr 2 (864)3, 24H 2 O, forms deep violet crystals, which behave when heated with water in a way similar to that mentioned above in the case of chromic chloride. The salt gives a green solution on boiling, which deposits a gummy mass on evaporation and from which barium chloride does not precipitate a sulphate. On standing for some time in cold water, the green salt changes and the violet form is deposited as crystals. Chrome alum is used as a source of chromic hydroxide in tanning and as a mordant. 776. The Chromates. The salts of chromic acid, which are used extensively as oxidizing agents and as pigments, are manu- factured from potassium chromate; the latter is made by roasting chromite FeO,Cr 2 Os, with potassium carbonate and lime in the presence of air. Under these conditions the iron in the mineral is oxidized to ferric oxide, Fe 2 O3, and the chromium to chromic anhydride, CrOs, which reacts with the carbonate to form potas- sium chromate, K 2 CrO4. The lime which is added to keep the CHROMIUM, MOLYBDENUM, TUNGSTEN, AND URANIUM 629 mixture porous and thus allow the air to act upon it, is also con- verted, in part, into calcium chromate. The product is treated with water to dissolve out the chromates, and enough potassium sulphate is added to precipitate the calcium as sulphate. Potassium chromate is an anhydrous salt and forms pale yellow crystals which are very soluble in water. Lead chromate, PbCrO 4 , has a brilliant yellow color and is the basis of the " chrome yellows " used as pigments, which are made by mixing it with other insoluble salts such as the sulphates of lead, barium, or cal- cium; it turns black in the presence of hydrogen sulphide. The solubilities of the chromates are very similar to those of the cor- responding sulphates. 777. Chromic Anhydride and the Dichromates. When con- centrated sulphuric acid is added to a strong solution of potassium chromate, chromic anhydride is precipitated in the form of dark red needles: K 2 CrO 4 + H 2 SO 4 = K 2 SO 4 + H 2 O + CrO 3 . The chromic acid, which is probably first formed, decomposes into its anhydride and water. The compound dissolves readily in water, but differs from sulphuric anhydride, SO 3 , in that but a small proportion of it unites with water to form an acid. It does unite, however, with neutral chromates to form stable compounds which are called dichromates: K 2 CrO 4 + Cr0 3 = K 2 Cr 2 O 7 . When the correct amount of sulphuric acid is added to a solu- tion of a chromate, a dichromate is formed; one-half of the salt ?s converted into potassium sulphate and chromic anhy- dride, and the rest unites with the latter to form the dichromate: K 2 CrO 4 + H 2 SO 4 = K 2 SO 4 + [CrO 3 ] + H 2 O K 2 CrO 4 + [CrO 3 ] = K 2 Cr 2 O 7 2K 2 CrO 4 + H 2 SO 4 = K 2 SO 4 + K 2 Cr 2 O 7 + H 2 O When potassium hydroxide is added to a solution of potassium dichromate it is converted into the chromate: K 2 Cr 2 O 7 = K 2 CrO 4 + [CrO 3 ] 2KOH + [CrO 3 ] = K 2 CrO 4 + H 2 O K 2 Cr 2 7 + 2KOH = 2K 2 Cr0 4 + H 2 630 INORGANIC CHEMISTRY FOR COLLEGES The change is evident on account of the fact that solutions of dichromates are red and those of chromates are light yellow. Potassium dichromate, which is orange-red in color, can be readily purified by crystallization from water on account of its relatively small solubility at room temperature (8 parts in 100 of water at 10). It was for this reason that it was formerly pre- pared in large quantities by the method described above (776) and was used in making other compounds of chromium. It has recently been replaced largely by the sodium salt, Na2Cr2O7,2H2O, which forms red crystals and is very soluble in water (109 parts in 100 at 15). Potassium dichromate ionizes largely to produce K + and Cr2<37~~ ions but the solution also contains free CrOa and CrO4~" ions formed as the result of the slight decomposition of the dichro- mate into the chromate and chromic anhydride. It is for this reason that if potassium dichromate is added to a solution of a salt the metal of which forms an insoluble chromate, the latter is precipitated : 2Ba(N0 3 ) 2 + K 2 Cr 2 O 7 + H 2 O = 2BaCr0 4 + 2KNO 3 + 2HNO 3 778. Ammonium dichromate, (NH^C^O?, is used in the sensitive material employed in making photographic prints by the carbon process. Paper is coated in the dark with a solution of gelatin and ammonium dichromate, in which lamp-black or other finely divided pigment is suspended, and allowed to dry. When it is exposed to light under a negative, a reaction takes place in the parts of the sensitive surface illuminated; the dichromate is reduced to chromium trioxide, which renders the gelatin in con- tact with it insoluble in warm water. The print is now washed carefully to remove the part of the sensitive surface which has not been acted upon by light, and the portion which has been developed is transferred to a suitable support. Powdered ammonium dichromate burns freely with a flame when it is ignited by means of a piece of paper which has been previously dipped into a solution of potassium nitrate and dried. The chromium changes in valence from 6 to 3, the oxygen liberated burns the ammonia to nitrogen, and the chromic oxide produced is formed as a light, voluminous, green powder. The reaction CHROMIUM, MOLYBDENUM, TUNGSTEN, AND URANIUM 631 which is a very striking one is represented by the following equa- tion: (NH 4 )2Cr 2 O7 = N 2 + 4H 2 O + Cr 2 O 3 . 779. The Use of Chromic Anhydride and Bichromates as Oxidizing Agents. The readiness with which chromic anhydride gives up a part of its oxygen is utilized extensively in effecting processes of oxidation. The compound chars paper and ignites alcohol at ordinary temperatures. In the presence of acids it is an active oxidizing agent in aqueous solutions. The acid facil- itates the process, since it unites with the trioxide formed as the result of the reduction of the anhydride: 2CrO 3 = [Cr 2 O 3 ] + 3O [Cr 2 O 3 ] + 3H 2 SO 4 = Cr 2 (SO 4 ) 3 + 3H 2 O 2CrO 3 + 3H 2 SO 4 = O 2 (SO 4 ) 3 + 3H 2 O + 30 Three oxygen atoms become available from 2 molecules of the anhydride if a substance is present that can be oxidized. For most oxidations, sodium or potassium dichromate and sulphuric acid are commonly used because the salts are cheaper than the anhy- dride. The acid first liberates the anhydride from the salt, and then the reactions indicated above take place. If the reaction K 2 Cr 2 O 7 + H 2 S0 4 = K 2 SO 4 + H 2 O + 2CrO 3 is added to these, the combination which expresses the reaction when the salt is used is as follows: K 2 Cr 2 7 + 4H 2 SO 4 = K 2 SO 4 + Cr 2 (SO 4 ) 3 + 4H 2 O + [30] This reaction occurs only when something is present that can be oxidized, and the fact is indicated by writing the formula for oxygen as 0; the equation which indicates the oxidation of ferrous sul- phate to ferric sulphate is as follows: 2FeSO 4 + H 2 SO 4 + O = Fe 2 (SO 4 ) 3 + H 2 O In order to combine this equation with the one given above, it must be multiplied by 3 so that the amount of oxygen available for the oxidation is equal to that taken up by the substance 632 INORGANIC CHEMISTRY FOR COLLEGES oxidized. When this has been done and the partial equations are added we have the following result : K 2 Cr 2 O 7 + 4H 2 SO 4 = K 2 S0 4 + Cr 2 (SO 4 ) 3 + 4H 2 O + [3O] 6FeSO 4 + 3H 2 S0 4 + [3O] = 3Fe 2 (SO 4 ) 3 + 3H 2 O K 2 Cr 2 O 7 + 6FeSO 4 + 7H 2 SO 4 = K 2 SO 4 + Cr 2 (SO 4 ) 3 +3Fe 2 (SO 4 ) 3 + 7H 2 O The reaction which takes place between potassium dichromate and hydrochloric acid can be written in a similar way : K 2 Cr 2 O 7 + 8HC1 = 2KC1 + 2CrCl 3 + 4H 2 O + [3O] 6HC1 + [3O] = 3H 2 O + 3C1 2 K 2 Cr 2 O 7 + 14HG1 = 2KC1 + 2CrCl 3 + 7H 2 O + 3C1 2 780. Reactions of Oxidation from the Point of View of Positive and Negative Valence. When mercuric chloride, HgCl 2 , is treated with stannous chloride, SnCl 2 , the latter is oxidized to stannic chloride, SnCl 4 , and the former reduced to mercury. In mercuric chloride the mercury has the valence +2, because it is combined with 2 atoms of a univalent negative element; in the metallic condition its valence is 0. The change in valence is 2, because 2, the original valence, 2=0, the final valence. The valence of tin in stannous chloride is +2 and in stannic chlo- ride +4, and, accordingly, as the result of the reaction the valence of tin has changed by +2, for +2, the original valence, +2 = +4, the final valence. The mercury " loses 2 valencies " and the tin takes these up, for it gains 2. As a consequence, 1 molecule of mercuric chloride will react with 1 molecule of stannous chloride: HgCl 2 + SnCl 2 = Hg + SnCl 4 . Since reactions of oxidation of this type involve the transfer of an element or a radical from one element to another, the loss in valence of the atoms of the element in the oxidizing agent must, evidently, equal numerically the gain in valence of the atoms of the element in the reducing agent, or, expressed in other words, the two changes in valence must be equal but of different sign. From the point of view of + and valence an element is oxidized CHROMIUM, MOLYBDENUM, TUNGSTEN, AND URANIUM 633 when there is an increase in its positive valence (or a decrease in its negative valence) ; and the reverse is true in reduction. In the case of the reduction of ferric chloride, Feds, by stan- nous chloride, the iron salt is reduced to ferrous chloride, FeCl 2 , and the change in valence of the metal is 1, for +3 1 = +2. The change of valence of the tin in stannous chloride is +2, con- sequently, 1 molecule of it will reduce 2 molecules of ferric chloride: 2FeCl 3 + SnCl 2 = 2FeCl 2 + SnCl 4 . The same principle applies when the negative valencies are involved in oxidation reactions. When ammonia is oxidized to nitric oxide the change in the valence of nitrogen is from 3 in ammonia to +2 in nitric oxide, NO. It will be recalled that the <-H electronic formulas of these two compounds are written N < H <-H and N 0, to indicate that the electrons involved pass from the more positive to the less positive element (442); and it will also be remembered that one so-called positive valence is established on an element for each electron lost and one negative valence for each electron gained. The total change in valence in oxidizing ammonia to nitric oxide is, therefore, from 3 to +2 which is +5, for 3 + 5 = +2. The change in valence of the element in the oxidizing agent will be, accordingly, 5. If free oxygen is the agent employed, it changes in valence from to 2 (H 2 O) and, consequently, 2J atoms will be required for each molecule of ammonia oxidized, 1NH 3 to 2? O, or expressed as a molecular ratio, 4NH 3 to 5O 2 . When ammonia is treated with chlorine it is oxidized to nitro- gen and hydrochloric acid is formed; in this case the change in valence of nitrogen is from -3 to 0, which is +3, and of chlorine from to 1; consequently, for each molecule of ammonia oxi- dized three atoms of chlorine will be required: 1NH 3 to 3C1 or 2NH 3 to 3C1 2 : 2NH 3 + 3C1 2 = N 2 + 6HC1. 781. The reactions of oxidation involving the use of potassium dichromate will next be considered from this point of view. In the salt the valence of chromium is 6, because it is derived from chromic acid, the anhydride of which is CrO 3 . Another way of determining the valence of chromium is to consider the positive 634 INORGANIC CHEMISTRY FOR COLLEGES and negative valencies of the atoms of which the salt is made up. In K2Cr2O? these are 2 potassium atoms each having the valence + 1, and 7 oxygen atoms each having the valence 2. In any compound the sum of the negative valencies equals the sum of the positive valencies, or, in other words, the sum of the valen- cies is numerically equal to (442) . Since in potassium dichromate there are 7 X 2 = 14 valencies due to the oxygen, there must be +14 valencies; of these the two potassium atoms furnish 2 and the rest, 12, must be furnished by the two chromium atoms, which, accordingly, have 6 each. When potassium dichromate acts as an oxidizing agent in the presence of acids, the chromium is reduced to the valence +3 and chromium salts are formed: the change in valence is from 6 to 3 or 3, and since each molecule of the salt contains 2 atoms of chromium the total change per molecule is 6. This amount of the dichromate will, accordingly, oxidize the amount of another substance which involves a change of + 6 in valence. If hydro- chloric acid is to be oxidized to chlorine the change in valence of chlorine is from 1 to 0, which is +1; consequently, 1 molecule of potassium dichromate will oxidize 6 molecules of hydrochloric acid to 6 atoms of chlorine. More of the acid will be required than this amount, however, because the chromium changes to the valence 3 in which it is a base-forming element and, con- sequently, is converted into chromic chloride, CrCla. Since 2 atoms of the element are involved, 6 molecules of hydrochloric acid will be required to furnish the chlorine for this purpose. And, further, the potassium in the salt is also converted into chloride and 2 molecules of the acid will be needed for this pur- pose, making a total of 6 + 6 + 2 = 14 molecules of the acid. The equation can now be readily written: K^C^O? + 14HC1 = 2KC1 + 2CrCl 3 + 3C1 2 + 7H 2 O. It should be noted that the oxygen in the dichromate is converted into water, the hydro- gen for the purpose coming from the acid. This fact makes it possible to see by inspection how many molecules of acid are required. In this case there are the 7 oxygen atoms, and the 14 atoms of hydrogen needed are furnished by 14 molecules of hydro- chloric acid. 782. If we consider next the oxidation of ferrous sulphate, FeSO4, to ferric sulphate, Fe 2 (864)3, by potassium dichromate in CHROMIUM, MOLYBDENUM. TUNGSTEN, AND URANIUM 635 the presence of sulphuric acid, the total change in valence for 1 molecule of the dichromate is 6. When iron changes from the ferrous condition in which the valence of the metal is 2 to the ferric condition in which it is 3 the change in valence is +1; as a con- sequence, 1 molecule of potassium dichromate will oxidize 6 molecules of ferrous sulphate. The number of molecules of sulphuric acid required is determined as shown above by the number of oxygen atoms in the oxidizing agent; the 14 atoms necessary to convert 7 atoms of oxygen into water are furnished by 7 molecules of sulphuric acid. The completed reaction is, accordingly, as follows: K 2 Cr 2 O 7 + 6FeS0 4 + 7H 2 S0 4 = K 2 S0 4 + Cr 2 (S0 4 ) 3 +3Fe 2 (S0 4 ) 3 + 7H 2 O As we have seen, when 1 molecule of potassium dichromate acts as an oxidizing agent in the presence of an acid, the valence change is 6. When oxygen gas acts as an oxidizing agent, the change in valence for each atom is 2, for the element has this valence in its compounds; consequently, 1 molecule of the dichromate is equivalent in oxidizing power to 3 atoms of oxygen. The examina- tion of reactions of oxidation from this point of view is often very helpful. For example, in oxidizing alcohol, C 2 H6O, to acetic acid, C 2 H 4 O 2 , 2 atoms of oxygen are necessary C 2 H 6 O + 2O = C 2 H 4 2 + H 2 O As a consequence, we see that 2 molecules of the dichromate, which are equivalent to 6 oxygen atoms, will be required to oxidize 3 molecules of alcohol which require 6 of oxygen to effect their oxidation. The equation for the reaction is as follows: K 2 Cr 2 O 7 + 3C 2 H 6 O + 4H 2 S0 4 = K 2 S0 4 + Cr 2 (S0 4 ) 3 + 3C 2 H 4 2 + 3H 2 O In this case but 4 molecules of the acid are required because only the potassium and chromium are converted into sulphates. 783. Analytical Reactions of Chromium. Solutions of chromic salts are bluish-violet or green in color (773) and yield, when treated with a soluble hydroxide, a bluish-green precipitate of the hydroxide, which is insoluble in ammonia. All the compounds of chromium when fused with sodium carbonate and potassium 636 INORGANIC CHEMISTRY FOR COLLEGES nitrate are converted into a chromate which is yellow; when fused with borax they give a green bead, which is not changed in color in the reducing flame. The soluble chromates give yellow and the dichromates red solutions; they are reduced by hydrogen sulphide, and insoluble chromic hydroxide and sulphur are formed. 784. Molybdenum. This element forms a number of oxides and chlorides, but its most important compound is molybdic anhy- dride, MoOs, from which a large number of salts have been made. The element is obtained chiefly from molybdenite, MoS2, by con- verting it, by roasting, into the trioxide, which is then reduced by heating it with hydrogen. Since the metal melts at a very high temperature it cannot be worked in the usual way; the powder obtained as the result of the reduction is pressed, and the ends of the blocks connected with terminals which furnish a powerful electric current. The heat generated by the passage of the current causes the particles to coalesce. Molybdenum is used in steel and is made into terminals for spark-plugs. To obtain it in the form of wire, the metal is drawn while red-hot through dies made of dia- mond set in steel. Ammonium molybdate, (NH^MoCU, which is made by dis- solving molybdic anhydride in ammonia, is used in the laboratory in testing for phosphates and in their quantitative determination. When a phosphate is added to a solution of the reagent in dilute nitric acid, and the mixture is heated, a yellow precipitate of ammonium phosphomolybdate is produced. The salt, which has the formula (NH 4 )3PO 4 ,12MoO3,6H 2 O, is insoluble in dilute acids, but dissolves in alkalies. 785. Tungsten. The element occurs as tungstates. When these are fused with sodium carbonate they are converted into sodium tungstate, Na2WC>4, which is soluble in water. Acids precipitate from the solution tungstic acid, EbWO^H^O, which yields tungsten on reduction with hydrogen. The metal is used extensively in making high-speed steel (748) and as a filament in electric-light bulbs. It is obtained in the form of wire in the way just described in the case of molybdenum. Tungsten melts at a higher temperature (3540) than any other element, and the dis- covery of a method to change into the form of a wire the powder formed as the result of the reduction of the oxide, revolutionized the electric-light industry. Carbon filaments were formerly used CHROMIUM, MOLYBDENUM, TUNGSTEN, AND URANIUM 637 in electric light bulbs, but when they were heated to a very high temperature the carbon vaporized, and depositing on the glass darkened it; in a short time the filament became disintegrated. On account of its high melting-point and low volatility, tungsten can be heated to a much higher temperature than carbon without acting in this way, and since the quantity of heat converted into light increases very rapidly with rise in temperature (455), the efficiency of a tungsten lamp is much greater than that of one in which the filament is made of carbon. A lamp of the latter type requires 3.25 watts per candle-power, whereas a tungsten lamp requires only 1.25 watts to produce the same amount of light. 786. Uranium. The element occurs in small quantities in several minerals which are complex in composition; in pitchblende it is present as an oxide, UsCU, associated with a large number of other elements; carnotite, which has the composition K2O,2UOs,- V2Os,3H2O, occurs in Colorado and is used as a source of radium which is found in minute amounts in uranium ores. The element forms a large number of compounds, the most important of which are derived from the oxide of the composition UOs, which is an acid anhydride and yields salts analogous in composition to the chro- mates and dichromates. Sodium uranate has the formula Na2UC>4, and sodium diuranate, the formula Na2U2O7,7H2O. The latter salt is used in coloring uranium glass, which exhibits the phe- nomenon of fluorescence, and is yellow or green depending on the way the glass is viewed. Uranium trioxide has weak base-forming properties; it forms basic salts of which uranyl nitrate, UO2(NO3)2, 7H2O and uranyl sulphate, UC^SCU, 3JH20, are examples. The radical which occurs in these salts, UC>2 ++ , is called uranyl; the name recalls that of sulphuryl chloride, SO 2 C1 2 (290). The uranyl salts are yellow and their solutions fluorescent, being yellow by reflected and green by transmitted light. Uranium and its compounds are of especial interest on account of the fact that they exhibit the phenomenon of radioactivity (799). EXERCISES 1. Compare the chemical behavior of (a) A1(OH) 3 , Fe(OH) 3 , andCr(OH) 3 , (6) Cr(OH) 2 and Fe(OH) 2 . 2. Write equations for all the reactions involved in making K 2 Cr 2 O 7 from chrome-iron ore. State why the process serves to separate the iron from the chromium. 638 INORGANIC CHEMISTRY FOR COLLEGES 3. Write ionic equations for the reactions which occur when a solution of Ba(NO 3 ) 2 is treated with one of K 2 Cr 2 O 7 . 4. Calculate the valence of the acid-forming element in the following compounds, using -f- and valencies: (a) CrPO 4 , (6) K 2 Mn 5 On, (c) KH 2 PO 2 , (d) CaSO 3 , (e) (NH 4 ) 2 Cr 3 O, . 5. Under certain conditions alcohol, C 2 H 6 0, is oxidized to aldehyde, C 2 H 4 O. Write the equation for the reaction which takes place when the oxidation is effected by means of K 2 Cr 2 O 7 in the presence of H 2 SO 4 . 6. Write an equation for the oxidation of H 2 S by K 2 Cr 2 O 7 as the result of which S and Cr(OH) 3 are formed. 7. How could you distinguish from each other the salts of the following metals: (a) Cr and Ni, (6) Cr and Fe, (c) Cr and Co, (d) How could you separate Fe(OH) 3 from Cr(OH) 3 ? 8. How could you obtain pure metallic iron and pure metallic chromium starting with a mixture of the sulphates of the metals? CHAPTER XLIV MANGANESE 787. We have just seen from a study of the chemistry of chromium that the valence of an element is the factor which largely determines the chemical behavior of its compounds. This important conclusion is even more strikingly illustrated in the case of manganese, which forms compounds in which the element shows the valencies 2, 3, 4, 6 and 7. In the elementary state manganese has the properties characteristic of metals and resem- bles closely iron. When it has the valence 2 it forms salts like those of magnesium and is a relatively strong base-forming ele- ment. With the valence 3, it functions as a weak base-forming element and yields salts which resemble those of aluminium al- though they are much less stable and are more readily hydrolyzed. With increase in the valence of the element to 4, compounds are produced which are like those of lead when it shows this valence; the oxide of the composition MnO2 has very weak acid-forming properties, and the corresponding chloride, MnCU, is unstable and like lead tetrachloride breaks down into the dichloride and chlorine. The element becomes strictly acid-forming in character when it has the valence 6; it forms salts, of which potassium manga- nate, K^MnCU, is an example, that are isomorphous with the sulphates. Its highest valence is 7, a fact which is in accord with its position in the seventh group in the periodic classification of the elements. With this valence manganese functions as a strong acid-forming element and yields salts called permanganates, which resemble the perchlorates in both physical and chemical properties; potassium permanganate, KMnO4, and potassium perchlorate, KC1O4, are examples of such salts. 788. Preparation and Properties of Manganese. The ele- ment occurs chiefly as oxides and hydrated oxides. Braunite, , manganite, MnO (OH) or M^Oa , H2O, hausmannite, 639 640 INORGANIC CHEMISTRY FOR COLLEGES and manganese spar, MnCOs, are important minerals. The chief ore of the element is pyrolusite, Mn(>2. Manganese may be obtained by reducing an oxide of the metal with aluminium, but the large amounts required in making iron alloys are made by reducing a mixture of an iron ore and a manganese ore in a blast furnace. Spiegeleisen and ferromanganese are made in this way; the former contains 10 to 20 per cent manganese, 4 to 5 per cent carbon, and the rest is iron; the latter contains 20 to 85 per cent manganese and 6 to 7 per cent carbon. Manganese resembles iron in appearance but its luster has a reddish tinge; it is hard and brittle, has the specific gravity 7.2 and melts at 1260. It oxidizes in the air, decomposes steam at red heat, and dissolves in acids to form manganous salts. When the element or any of its oxides is heated in the air the final product has the composition MnaO^ 789. Manganous Compounds . Manganous hydroxide, Mn(OH)2, is precipitated when a base is added to a solution of a manganous salt; when pure it is white, but it is soon oxidized by the air more or less completely to manganic hydroxide, Mn(OH)3, which is brown; the reaction is like that which takes place in the case of ferrous hydroxide. The hydroxide, like ferrous and magnesium hydroxides, is not precipitated by ammonium hydrox- ide in the presence of ammonium salts (589). Manganous chloride, MnCl2,4H2O, is obtained by treating manganese dioxide with hydrochloric acid; the salt crystallizes from water, is pale pink in color, and is stable in the air. Man- ganous sulphate, MnS(>4, forms a number of hydrates; the one obtained by evaporation of a solution at room temperature is a pentahydrate, which is isomorphous with copper sulphate. Manganous carbonate, MnCOs, is white when pure, but is usually colored brown owing to the presence of manganic hydroxide which is formed as the result of the oxidation of the small amount of manganous hydroxide produced as the result of the partial hydrolysis of the carbonate. It is decomposed at a compara- tively low temperature by heat, and for this reason is used under certain conditions in the laboratory as a source of carbon dioxide. Manganous sulphide, MnS, is formed by the action of ammonium sulphide on a manganous salt; it is a flesh-colored amorphous precipitate, which, like ferrous sulphide, is soluble in dilute acids. MANGANESE 641 All the soluble manganous salts have a pink color; those derived from the strong acids are but slightly hydrolyzed and are stable in the air, and in the latter respect differ from the salts of stannous tin and ferrous iron. 790. Manganic Compounds. Manganic hydroxide, Mn(OH)s, is formed as the result of the hydrolysis of manganic sulphate, Mn2(SO4)3, and by the oxidation of the solution formed by dissolving manganous hydroxide in solutions of ammonium salts. It is brown-black in color and dissolves in hydrochloric acid to form manganic chloride, MnCls, which readily decomposes into manganous chloride and chlorine. Like other unstable com- pounds of this kind, it forms double salts which are stable; for example, a compound of the formula MnCls, 2KC1 can be pre- pared by dissolving manganese dioxide in a solution containing hydrochloric acid and potassium chloride. The marked difference between iron and manganese is that in the case of the former the salts containing the element in the trivalent condition are stable in the air, whereas in the case of manganese salts, the salts of the metal in the bivalent condition are stable. In the trivalent condition manganese is a less active base-forming element than iron, and its salts are more com- pletely hydrolyzed than those of ferric iron. 791. Manganites. Manganese dioxide, MnO2, like silicon dioxide, SiO2, acts as an acid anhydride when fused with alkalies; like other anhydrides of this class, in which acid-forming proper- ties are developed to but a slight degree, it yields salts that are complex in composition and contain a number of molecules of the anhydride to one molecule of the basic oxide. There are many silicates, chromates, tungstates, and molybdates of this type that are derived from anhydrides which are very weakly acidic. When manganese dioxide is fused with potassium hydroxide, a number of manganites are formed, one of which has the formula K 2 O,5MnO 2 . When calcium hydroxide is added to a solution of a manganese salt and air is blown through the mixture, the man- ganous hydroxide precipitated is, in part, oxidized in the pres- ence of the base, and manganites of the formulas CaO,MnO2 and CaO,2MnO 2 are produced. A part of the manganese remains in the bivalent condition and unites with some of the dioxide to form manganous manganite, MnO,MnO 2 . These reactions are 642 INORGANIC CHEMISTRY FOR COLLEGES utilized in recovering the manganese used in the process for the manufacture of chlorine from manganese dioxide and hydro- chloric acid. The so-called Weldon " mud " obtained as out- lined above is treated with hydrochloric acid, and chlorine is obtained. 792. Manganates. When any compound of manganese is fused with potassium carbonate and potassium nitrate, the melt becomes green in color and potassium manganate, K 2 MnO 4 , is obtained: MnO 2 + K 2 CO 3 + O = K 2 MnO 4 + C0 2 . The salt crystallizes from water in anhydrous greenish-black crystals, which give a deep green solution when dissolved in water. Manganese can have the valence 6 only when it is in com- bination with a base-forming element. If an acid is added to a solution of a manganate, the manganic acid, H 2 MnO 4 , liberated decomposes at once into permanganic acid, HMnO 4 , and man- ganese dioxide. The change can be readily observed because the manganates are green and permanganic acid and its salts are red. When a manganate is dissolved in water it hydrolyzes to some extent: K 2 MnO 4 + 2H 2 ^ 2KOH + H 2 MnO 4 The manganic acid formed then decomposes: 3H 2 Mn0 4 = MnO 2 + 2HMnO 4 + 2H 2 The permanganic acid next interacts with a part of the alkali formed as the result of the hydrolysis of the manganate : KOH + HMnO 4 = KMnO 4 + H 2 O The combination of these partial equations leads to the following equation: 3K 2 MnO 4 + 2H 2 O = 2KMnO 4 + 4KOH + MnO 2 Since potassium manganate is stable in the presence of alkalies, its hydrolysis takes place completely only if an acid is added to neutralize the hydroxide formed as the result of the hydrolysis. It is for this reason that carbon dioxide is passed into the solution in effecting the change of manganate to permanganate, and when MANGANESE 643 the equation for this reaction is combined with the one just given we arrive at the final equation for the reaction : 3K 2 MnO 4 + 2CO 2 = 2KMnO 4 + 2K 2 CO 3 + MnO 2 793. Permanganates. Potassium permanganate was formerly prepared by making use of the reactions just discussed, but since the process involved the conversion of one-third of the man- ganese into the dioxide, the salt is now manufactured by oxidizing a solution of potassium manganate by means of chlorine : 2K 2 MnO 4 + C1 2 = 2KMnO 4 + 2KC1 It is also prepared by direct oxidation in an electrolytic cell: 2K 2 MnO 4 + O + H 2 O = 2KMnO 4 + 2KOH The permanganate separates as crystals near the anode. The hydroxide formed collects at the cathode, and is used in converting more manganese dioxide into manganate by the fusion process. Potassium permanganate dissolves in cold water (1 part in 16) and forms purple crystals. It is an active oxidizing agent, which can ignite certain organic compounds; when a drop of glycerine is placed on the powdered salt, the former burns with a blue flame. Potassium permanganate is used as a disinfectant, in bleaching and dyeing, and in medicine. Sodium permanganate was used during the recent war in gas masks as a protection against arsine and other oxidizable gases, and was incorporated into granules containing sodium hydroxide, lime, and cement. The bases were used to decompose halogen compounds and the cement to render the mass porous. Permanganic acid can be obtained in the form of hydrated crystals by adding sulphuric acid to a solution of barium perman- ganate, filtering off the barium sulphate, and evaporating the fil- trate at a low temperature. The acid decomposes at 32 into manganese dioxide, water, and oxygen. It has also been prepared by electrolyzing a solution of potassium permanganate. The "acid collects at the anode, which is placed in a porous cup, and potassium hydroxide at the cathode. When potassium permanganate is cautiously treated with concentrated sulphuric acid, an oil, which gives a violet vapor, is 644 INORGANIC CHEMISTRY FOR COLLEGES formed; it has the composition Mi^O?, and is the anhydride of permanganic acid. The oil is explosive and sets fire to paper and other organic substances. 794. Potassium Permanganate as an Oxidizing Agent. The fact that potassium permanganate is such an active oxidizing agent, makes it a very valuable reagent in volumetric quantitative analysis. The solution of the salt is stable provided it is kept free from any oxides of manganese, which catalyze its reduction to manganese dioxide. Since it slowly oxidizes paper and other organic substances, it must be protected from these as well as from the dust of the air, which contains organic material. No indicator is needed when we are using the solution as an oxidizing agent, for its purple color disappears when it has been reduced. When potassium permanganate oxidizes in the presence of acids it is reduced to a manganous salt and the change in valence is from +7 to +2. If ferrous sulphate is being oxidized, for example, to ferric sulphate, the change in valence of the iron is from +2 to +3; as a consequence, 1 molecule of the perman- ganate will oxidize 5 molecules of ferrous sulphate. When the oxidation occurs in neutral or alkaline solutions the permanganate is reduced only to manganese dioxide and the valence change is from +7 to +4. It is important to remember that in acid solution the salt furnishes oxygen equivalent to 5 " valencies " or equivalents (2J atoms of oxygen), and in neutral or alkaline solutions, 3 equivalents. With these facts in mind it can be seen readily that 1 molecule of potassium permanganate will oxidize in acid solution 5 molecules of a chromous" salt to a chromic 111 salt, 2J molecules of a stannous 11 salt to a stannic iv salt, 2| molecules of sulphur iv dioxide to sulphuric vi acid, 5 molecules of hydriodic" 1 acid to iodine , 2J molecules of nitrous 111 acid to nitric v acid, etc. The equations for the reactions can be written by applying the method described at length when chromic acid was considered (779). The case of the oxidation of ferrous sulphate in the presence of sulphuric acid may be taken as an example. One molecule of the permanganate will oxidize 5 molecules of the salt, since the change in valence of the manganese is 5 and that of the iron salt is 1. Since ferric sulphate contains 2 atoms of iron, we must take an even number of molecules of ferrous sulphate 10 MANGANESE 645 molecules of the latter will make 5 of ferric sulphate. We require for this amount, evidently, 2 molecules of potassium permanganate. These contain 8 oxygen atoms which are converted into water in the reaction, and, consequently, 8 molecules of sulphuric acid are necessary to furnish the required 16 hydrogen atoms. The completed reaction is then as follows: 2KMnO 4 + 10FeSO 4 + 8H 2 SO 4 = K 2 SO 4 + 2MnSO 4 + 5Fe 2 (SO 4 ) 3 + 8H 2 O 795. Analytical Reactions of Manganese. Compounds of the element are converted into a green manganate when fused with potassium carbonate and potassium nitrate. When treated with ammonium sulphide manganous salts give a flesh-colored precipitate of manganous sulphide, which dissolves in dilute acids. Manganates are green and permanganates are violet; they are converted by ferrous salts, zinc and sulphuric acid, and by other reducing agents into manganous salts, which are very light pink or practically colorless in dilute solutions. EXERCISES 1. How could Mn2O 3 be distinguished from Fe?O 3 through the use of HC1? 2. Complete and balance the following: (a) FeSO 4 -f K 2 MnO 4 -f- H 2 SO 4 = , (b) SnCl 2 -f K 2 MnO 4 + HC1 = , (c) KMnO 4 + HC1 = , (d) KMnO 4 + S0 2 + H.SO 4 = , (e) KMnO 4 + KNO 2 + H 2 SO 4 = . 3. The strength of a solution of KMnO 4 can be determined by titratimg it with Na 2 S 2 O 3 after treatment with KI and dilute sulphuric acid. Write equations for all the reactions involved. 4. Write equations for the reactions that take place when an alloy of Fe, Mn, and Cr is dissolved in hot aqua regia. 5. Explain why Fe can be separated by the basic acetate method (756), from Mn and not from Cr. 6. (a) Write an equation for the reaction by which alcohol, C 2 H 6 O, is oxidized to CO 2 + H 2 O by KMnO 4 in the presence of H 2 SO 4 . (6) Write an equation for the oxidation if no acid is present. CHAPTER XLV RADIOACTIVITY. THE STRUCTURE OF ATOMS 796. The discovery of X-rays and of radium (652) brought to light a new type of energy the study of which has resulted in extending to a marked degree our knowledge of matter and its transformation. It is possible to count individual atoms and determine their weight, and the facts which have been discovered lead to the reasonable conclusion that the atoms of the elements are made up of simpler units, which appear to be hydrogen and helium. Before the facts upon which these views are based can be given, some of the phenomena concerned must be described briefly. 797. X-Rays. A tube to produce X-rays is made by sealing into a glass bulb two metallic conductors, and exhausting the bulb to the highest vacuum obtainable. When an electric cur- rent from an induction coil is sent through the tube, a stream of electrons, which are charges of negative electricity, is given off from the cathode. When these fall upon any substance placed opposite the cathode (the anti-cathode), vibrations, called X-rays, are produced which are of the same nature as those of visible light. They differ from the latter only in wave-length. The wave-lengths of visible light are of the order of 10~ 5 cm., while those of the X-rays are about one-thousandth as large (10~ 8 cm.). X-rays will pass through materials through which the longer light rays will not pass; it is this property of the rays which is utilized in X-ray photography. 798. Atomic Numbers. When white light falls upon a " grat- ing " made by ruling lines on a polished metallic surface, the light is broken up into its constituents and a spectrum is formed. In order to obtain this effect, the distances between the unscratched surfaces of the grating which reflect light must be of the order of the wave-length of light. Gratings have been made by ruling 646 RADIOACTIVITY. THE STRUCTURE OF ATOMS 647 200,000 lines to the inch. X-Rays of different wave-lengths are separated in a similar way, when they fall on a surface which is so constituted that alternate parts will reflect the waves. It is impossible to make a grating of this kind mechanically, because the distances between the reflecting surfaces must be of the order of the wave-lengths of the rays. It was discovered by Laue, in 1912, that when X-rays fall upon a crystal of salt, the surface of the crystal acts as a grating as a result of the orderly arrangement of its atoms, which serve as the reflecting surface. In this way it is possible to determine the wave lengths of the X-rays given off from any substance when it is the anti-cathode in an X-ray tube. The spectrum produced is not visible, as it is in the case of light, but it can be photographed in the ordinary way, because X-rays act upon a photographic plate. Moseley in 1914 examined the X-rays produced when different elements served as the anti-cathode in an X-ray tube. He found that each substance produced waves of a definite wave-length, which are characteristic of the element. And what was more striking, he found that there was a direct relation between the atomic weight of the element and the wave-length of the X-rays produced. As the atomic weight increased, the wave-length increased. When the elements were arranged according to the wave-lengths of the X-rays they gave off, they followed in the order in which they occur in the periodic classification. The exceptions which are met with when the elements are arranged according to their atomic weights, disappeared, however. It will be recalled that the atomic weight of tellurium places it in the periodic classification in the halogen family, and that of iodine in the sulphur family, and the order in which iron, cobalt, and nickel fall is anomalous. When these elements are arranged according to the wave-lengths of their characteristic X-rays, they fall in the order demanded by their chemical properties. It is thus seen that the atomic weight of an element, which is a measure of its mass, is not its most fundamental property. The X-rays emitted by an atom when it is struck by an electron, are produced as the result of the vibrations set up in the atom. These are determined, in all probability, by the electrical charges on the atom. In classifying the elements according to the wave-lengths of 648 INORGANIC CHEMISTRY FOR COLLEGES the X-rays they emit, a number is assigned to each element accord- ing to its position in the series. Hydrogen is numbered 1, helium 2, lithium 3, etc. The number assigned to each element is called its atomic number. The number of each element is given in the table illustrating the periodic classification, which is printed on the page facing the back cover of this book. It will be seen that a more exact statement of the properties of elements than that enunciated by Mendelejeff, can now be given; it is the prop- erties of the elements are a periodic function of their atomic num- bers. Mendelejeff 's generalization was based on the mass of the atoms, whereas the newer classification is based on energy rela- tionships, which appear to be more fundamental. An interpretation of the atomic numbers leads to the con- clusion that there are 92 elements, since 92 is the atomic number of uranium, and we have no evidence that heavier elements exist. Of these, 87 have been discovered. The places left vacant in the table and indicated by dots, . . . . , are those into which the five undiscovered elements should be put; their atomic numbers are 43, 61, 75, 85, and 87, respectively. 799. Radioactivity. The salts of radium constantly emit three kinds of " rays " which differ markedly from one another. The so-called a-rays have been shown to be made up of positively charged helium atoms, which are produced as the result of the loss of two electrons by each helium atom. It will be recalled that an electron is a negative charge of electricity, which has an appar- ent mass equal to about 18 1 00 that of a hydrogen atom. The charged helium atoms are called a-particles; they each carry the charge equal to that on a bivalent metallic ion, and when they lose this charge the helium has the properties of the gas extracted from natural gas and other sources (331). The /3-rays consist of electrons which move with about the same velocity as that of light. The 7-rays are X-rays produced as the result of the bombardment of the atoms of radium by the electrons which are given off by the element. During the disintegration of radium a gas is produced which was originally called radium emanation; it was shown to be an element that resembled argon in properties, and was later named niton. A determination of the density of the gas gave results that led to the atomic weight 222.4. The atomic weight of niton RADIOACTIVITY. THE STRUCTURE OF ATOMS 649 places it in group of the periodic classification of the elements; and this position is in accord with its properties. The fact that uranium is radioactive led to the view that the radium which is associated with it was probably produced from uranium, and an investigation of the subject showed that the view was correct. A sample of a uranium salt which had been freed from radium was found after about three weeks to contain the latter, and in about one year the relation between the amounts of the two elements present became constant (1 gram of uranium to 3.2X10" 7 grams of radium). This relationship makes clear why such small quantities of radium are obtained from pitchblende (652); from 1 ton of the mineral about 0.2 gram of radium can be extracted. 800. Radium is not the first product formed from uranium; the latter decomposes into helium and an element called uranium-Xi, which through the loss of electrons only passes into uranium-X2; this gives U2, which yields helium and ionium, and the latter then changes on disintegration into helium and radium. The decomposition of radium, as we have seen, produces helium and niton. The radio-active transformations continue through Ra-A, Ra-B, Ra-C, Ra-Ci, Ra-D, Ra-E Ra-F, and, finally, lead is produced. When helium, which has the atomic weight 4, is given off in the transformations, the atomic weight of the element formed is 4 less than that of the element from which it was produced. In the complete change from uranium, which has the atomic weight 238.2, to lead, there are eight radioactive transformations which produce helium; as a consequence, the atomic weight of lead should be 8 X 4 = 32 less than the atomic weight of uranium, that is, 206.2. Since the atomic weight of lead is 207.2, it seemed of importance to study the lead that was associated with uranium in its ores, for the metal formed under these conditions should have a lower atomic weight than 207.2, if the view is correct that it is the final disintegration product of uranium. Lead was isolated from uranium ores and after it had been freed from all other elements its atomic weight was determined. The metal obtained from different radioactive sources gave values for the atomic weight which varied from 206.4 to 206.8. The compounds prepared from the lead which has the atomic weight 650 INORGANIC CHEMISTRY FOR COLLEGES 206.4 possess the identical physical and chemical properties ex- hibited by the analogous compounds prepared from ordinary lead, which has the atomic weight 207.2. These results lead to the con- clusion that the atomic weight of an element is not its most charac- teristic property. The two forms of lead are called isotopes, the name signifying that the two varieties have different atomic weights, but that in all other properties they are identical. Since the announcement of this important discovery in the case of lead, many attempts have been made to isolate the isotopes of other elements, but adequate experimental evidence has not been furnished in any case. The fact that isotopes of other elements exist, has been demonstrated, however. On account of the fact that the atomic weight of lead from sources which are not radioactive is 207.2, it appears probable that lead of this kind is a mixture of isotopes, one of which is derived from uranium and the other of which possesses an atomic weight greater than 207.2. This conclusion receives striking con- firmation from the fact that radioactive thorium, which has the atomic weight 232.4, is the first member of a series of radioactive elements, which also pass finally into lead; the atomic weight of the latter in this case is 208.4, because the changes involve the loss of 6 helium atoms: 232.4 -6X4 = 208.4. A determina- tion of the atomic weight of a sample of lead obtained from a thorium mineral gave the value 208.4. It seems highly prob- able, therefore, that ordinary lead is a mixture of two isotopes, one of which is derived from uranium and the other from thorium. 801. Radioactive changes differ from ordinary chemical transformations in that the rate at which they take place does not vary with temperature. The chemist knows how to modify the reactions that occur between atoms, but up to the present no one has been able to control radioactive disintegrations, which result from the spontaneous decomposition of the atoms them- selves. The rate at which radium and other radioactive elements decompose is proportional to the amount of the substance under- going the change; consequently, as the amount diminishes the rate gets slower and slower. In order to express the rate at which the several elements of this class disintegrate it is usual RADIOACTIVITY. THE STRUCTURE OF ATOMS 651 to state the time required for one-half of any given amount to decompose. In the case of radium this time is 1850 years, and in the case of niton 3.85 days. One atom of radium out of each 100,000,000,000 undergoes decomposition per second. 802. There is a marked difference in the " life " of other ele- ments of this kind; radium-C exists for a fraction of a second, whereas in the case of uranium, the period for one-half decom- position is approximately 5 X 10 9 years. It has been deter- mined that 1 gram of uranium gives off helium at the rate of 1 c.c. in 16,000,000 years. The mineral fergusonite contains 26 c.c. of helium for each gram of uranium present. If the gas -was produced as the result of the disintegration of uranium, the process must have been taking place for 416 million years. The amount of energy set free as heat when radioactive trans- formations take place is very great; it has been calculated that the total energy liberated when an atomic weight of niton changes to radium-A is equal to 60,000 horse-power for one day. 803. Avogadro's Number. The a-particles expelled from radium travel at a velocity of about 0.1 that of light, and when they strike certain substances, zinc sulphide, for example, they render the latter luminescent. The instrument used to observe this phenomenon is called a spinthariscope. It consists of a small tube at one end of which is placed a trace of a radium salt in front of a screen covered with zinc sulphide, and at the other end of which is a lens through which the scintillations produced can be seen. It is possible to determine the number of helium atoms given off by radium in a given time by using a very small amount of the salt and placing the screen at a considerable distance from it. In this way the number of scintillations can be reduced to such an extent that they can be readily counted. It was determined that 1 gram of radium gives off 1.36 X 10 11 a-par- ticles per second, or 4.29 X 10 18 helium atoms per year. The helium given off by radium has been collected and measured, and it was found that the gas was formed at a rate of 0.156 c.c. per year. It follows, therefore, that 0.156 c.c. of helium contains 4.29 X 10 18 atoms, from which can be calculated that 1 c.c. of the gas contains 2.75 X 10 19 atoms and 1 gram-atomic-weight 6.16 X 10 23 atoms. Since helium is a monatomic gas, the num- ber of atoms present in its gram-atomic-weight is equal to the 652 INORGANIC CHEMISTRY FOR COLLEGES number of molecules present in the gram-molecular-weight of any other gas; 6.16 X 10 23 is, therefore, the number of molecules in a gram-molecular-weight of a gas (22.4 liters). This number has been arrived at in a variety of independent ways, and there is, therefore, strong evidence of its correctness; it is known as Avoga- dro's number. 804. The Structure of the Atom. We have seen that the radioactive elements are constantly undergoing disintegration as a result of which negative charges of electricity are produced. It has also been pointed out (565) that when a circuit made up of two metallic wires is heated at one of the junctions, an electric current is set up as the result of the flow of electrons through the wires. When a metallic wire is highly heated it gives off electrons into space, and if the wire is placed in a vacuum tube which con- tains also a sheet of metal that is charged positively, the negative charges pass through the vacuum from the wire to the sheet of metal and an electric current is produced. Such facts as the above lead to the conclusion that the atoms of which matter is composed are made up of electrical charges / which become evident when the atom is heated or when it spontaneously decom- poses. These views have been given concrete expression by the scientists who have studied the phenomenon outlined above, and have led to a conception of the structure of atoms which is the guiding principle in the research of to-day. Among the inves- tigators who have carried out the experimental work and have formulated the theories are J. J. Thompson, Rutherford, Soddy, and Ramsay. The discoveries which led to the active investiga- tions in this field were made by Crookes (cathode rays in 1878), Rontgen (X-rays, 1895), Becquerel (radioactivity of uranium compounds, 1896), and M. and Mme. Curie (radium, 1903). A large number of physicists and chemists have aided in developing the conception of atomic structure given below. 805. An atom is thought to be made up of a central nucleus composed of positive charges of electricity surrounded by a number of negative charges, placed at a distance from the nucleus, which is very small compared with the diameter of the space occupied by the atom. The atoms differ from one another in the number of positive and negative charges which they contain. The weight RADIOACTIVITY. THE STRUCTURE OF ATOMS 653 of the atom is associated with the positive charge. The negative charges are arranged around the nucleus in one or more shells, and those making up the outer shell are capable of transference to another atom; these are the so-called valence electrons, which function when chemical union between two atoms takes place. This general conception is applied to the several elements as follows: The hydrogen atom is considered to be made up of one positive charge and one electron, which is in vibratory motion and can function as a valence electron. Helium, the second element, consists of a positive nucleus and two electrons which are so firmly bound that they cannot be transferred to other atoms; the element is inert chemically. The succeeding elements are made of an inner core composed of a positively charged nucleus and electrons which do not function in chemical transformations, and an outer sphere of valence electrons, the number of which varies progressively as we pass from one element to the next in a series in the periodic classification. Lithium has 1 valence electron, beryllium 2, boron 3, carbon 4, nitrogen 5, oxygen 6, fluorine 7. Neon, the next element, has 8, and since it is inert, the conclusion is drawn that this arrangement leads to such a compensation of the forces in the atom that no electron can leave it. In the building up of the next element, sodium, by the addition of positive charges and electrons, those already present and those added, pass to the inner core except one electron which resides on a new outer shell. As we pass through the series as before, the next element, magnesium, has 2 valence electrons, aluminium 3, and so on to chlorine, which has 7, and to argon with 8, which, like neon, with 8 electrons in the outer shell, is inactive. The building up is continued through the elements in this way, and as a result there is a recurrence of elements with from 1 to 7 valence electrons. The elements thus exhibit a periodicity in the number of valence electrons they possess. The position of the electrons in the outer sphere is indicated on page 654. In each case, however, the inner core of the atom, which is not represented, is different. According to the electronic conception of the atom, when two elements combine, a valence electron passes from one element to the other. When,, for example, sodium chloride is formed, the metal loses its electron, which takes up its place in the outer sphere of the chlorine atom; the two atoms are held together by the 654 INORGANIC CHEMISTRY FOR COLLEGES electrical field of force set up between them as the result of the transfer. When the resulting compound is dissolved in water and the atoms are separated, the sodium having lost an electron, which is a negative charge, becomes a positively charged ion; and the chlorine having received an additional negative charge becomes a negative ion. FIG. 44. 806. Combination between atoms can take place in a different way from that just described. The two atoms may share elec- trons so that neither loses any. When 2 fluorine atoms unite to form a molecule, F2, it is assumed, according to this theory, that the atoms so arrange themselves that they share 2 electrons in common in such a way that two groups of 8 electrons are formed. This is represented in Fig. 45. Each fluorine atom has 7 elec- trons, and the arrangement of the two atoms indicated has 14 electrons; two electrons are shared in common by the two atoms. Compounds of this kind do not have the field of force between the two atoms that is set up in the way indicated in the case of sodium chloride; they are not ionized and are very stable. There appears to be a tendency for atoms to arrange themselves so that 8 electrons are present in the outer sphere; and, as we have seen, this configuration is the most stable and is present in the inert FIG. 45. RADIOACTIVITY. THE STRUCTURE OF ATOMS 655 The positive valence that an atom can show is determined by the number of electrons it has in its outer sphere, because when an element functions as a positive element it loses electrons and, therefore, assumes a positive charge. The negative valence is determined by the difference between 8 and the number of valence electrons, for the maximum number of the latter that can be present in the outer sphere is 8. If, for example, an element possesses 7, as in the case of chlorine, it can take up but 1 electron; its negative valence is, accordingly 1. On account of the signifi- cance of the number eight this theory is called the octet theory of valence. It has been developed by G. N. Lewis and also by Langmuir, who has used it to correlate many chemical facts. 807. It will be seen from the above that in ordinary chemical phenomena the outer sphere of the atom is alone concerned. The inner core maintains its identity in all such transformations, and the elements as we know them pass from one compound to another without change. In radioactive transformations, the inner core of the atom undergoes change and, as a result new ele- ments are formed. Since helium is given off in many of these changes, we must conclude that it is a constituent of all these atoms, and the reasonable assumption is that all atoms, except hydrogen, contain helium. Since the atom of helium weighs 4 and the atomic weights of the elements are not multiples of this number, we must conclude, further, that a lighter atom must be present in the complex atom of many elements. The most rea- sonable view to take is that hydrogen is the other element involved. Inspired by this hypothesis Rutherford recently subjected pure nitrogen to a bombardment by a-particles and obtained as a result a trace of hydrogen. No hydrogen was obtained from oxy- gen. These results are significant in view of the fact that the atomic weight of oxygen, 16, is a multiple of 4, whereas that of nitrogen, 14, could not be made up solely of helium. It is worthy of note that the elements which are sponta- neously undergoing radioactive disintegration are those of high atomic weight the ones that contain a large number of the simpler units of which the elements are composed. It is possible, there- fore, that the decomposition of the heavier atoms that do not show radioactivity may be effected more readily than that of the light elements, when the means to effect such changes are at hand. 656 INORGANIC CHEMISTRY FOR COLLEGES 808. General Conclusions. From the very brief and ele- mentary account of the investigations of the phenomena of radio- activity which has been given, it will be seen that it is now pos- sible to study the complex nature of the atoms, which up to recent years have been the final chemical units of matter. The con- ception of these units has changed as matter has been studied, and as the means by which matter can be decomposed have been learned. When heat energy alone was available, substances that are not decomposed by heat at the temperature obtained by chem- ical reactions, were thought to be elementary. When electrical energy was discovered and it was learned how to make and apply it, many substances undecomposed before were shown to be com- plex. Lime, for example, yielded a new metal calcium. And now, the energy developed when radioactive transformations take place carries us one step farther; when it is applied to the atom that resists ordinary electrical energy, the atom yields and decomposition takes place. Electricity was known a long time before it was found out how to generate it from other kinds of energy and how to use it to effect the decomposition of chemical compounds. It is not unreasonable to think that a way will be found to control or per- haps produce radioactivity, and when this happens practical transmutation of the elements may be possible. The elements are a storehouse of uncalculable energy, and if a way can be found to bring about their decomposition this energy can be made available for the uses of man. The study of radioactivity has led not only to the view that atoms are complex, but has shown that one of the fundamental laws upon which the quantitative study of chemical phenomena is based, is not an exact statement of the truth. The law of def- inite proportions does not apply to isotopes; the compounds made from lead obtained from radioactive sources do not contain a fixed proportion of the elements present. In one case 208.4 grams of lead unite with 2X35.45 grams of chlorine to form lead chloride and in the other case 206.4 grams of lead unite with the same weight of the halogen. But the newer knowledge does not in any way modify the facts which have been learned as the result of the intensive study- of chemical phenomena; they teach us that there is much ahead to be discovered and that we must mod- RADIOACTIVITY. THE STRUCTURE OF ATOMS 657 ify our conclusions and theories as the science grows. The greater the modifications necessary, the greater the growth of the science. The last decade has advanced our knowledge of matter to such a remarkable degree that it is impossible to foresee what lies just ahead. APPENDIX I. TEMPERATURE SCALES On the centigrade scale, the freezing-point and boiling-point of water are called, respectively, and 100. On the Fahrenheit scale, these points are called, respectively, 32 and 212. Therefore, 100 degrees C. = 212 32 = 180 degrees F. One degree C. = f degree F. and 1 degree F. = f degree C. Temperatures on one scale can be changed to temperatures on the other by the use of the following formulas, in which C is the reading on the centigrade thermometer and F is the reading on the Fahrenheit thermometer. C = f (F - 32), F = f (C) + 32 Absolute temperature, T, = C+ 273. II. METRIC SYSTEM Length. 1000 Millimeters = 100 centimeters = 1 meter = 39. 37 inches. 1 Inch = 2.54 centimeters. Volume. 1000 Cubic centimeters = 1 liter = 0.03532 cu. ft. = 61.03 cu. in. = 1.057 quarts = 34.1 fluid ounces. 1 Fluid ounce = 29.57 c.c. 1 Quart = 0.9434 liter. 1 Cubic foot = 28.32 liters. Weight. 1000 milligrams = 100 centigrams = 10 decigrams = 1 gram = 0.03527 ounce = 15.43 grains. 1 Kilogram = 1000 grams = 2.205 pounds avoirdupois. 1 Pound avoirdupois = 453.6 grams. 1 Ounce avoirdupois = 28.35 grams. 1 Metric ton = 1000 kilos = 2205 Ibs. III. ENERGY UNITS 1 Calorie (cal.) = heat to raise the temperature of 1 gram of water 1 degree centigrade (at 15). 1000 Calories = 1 large calorie (Cal.). 1 British thermal unit (B.t.u.) = heat required to raise the temperature of 1 Ib. of water 1 degree Fahrenheit at its maximum density = 252 cal. 659 660 APPENDIX Joules = coulombs X volts. 1 Joule = 0.238 cal. 1 Calorie = 4.187 joules. 1 B.t.u. = 1055 joules. IWatt = 1 joule per second. 1 Horse-power = 33,000 ft.-lbs. per minute = 746 watts. IV. MOLECULAR VOLUMES 1 Gram-molecular-weight of a gas occupies 22.4 liters at and 760 mm. 1 Pound-molecular-weight of a gas occupies 359 cubic feet at and 760 mm. V. VAPOR PRESSURE OF WATER The temperatures are given in degrees centigrade, and the pressure in millimeters of mercury. Temp. Pressure. Temp. Pressure. Temp. Pressure. 4.6 16 13.5 26 25.1 5 6.5 17 14.4 27 26.5 8 8.0 18 15.4 28 28.1 9 8.6 19 16.3 29 29.8 10 9.2 20 17.4 30 31.5 11 9.8 21 18.5 31 33.4 12 10.5 22 19.7 32 35.4 13 11.2 23 20.9 33 37.4 14 11.9 24 22.2 34 39.6 15 12.7 25 23.6 35 41.8 VI. GROUPS OP THE COMMON METALS IN QUALITATIVE ANALYSIS Group 1. Precipitated by HC1 : AgCl, HgCl, PbCl 2 . Group 2. Precipitated by H 2 S in HC1 solution: HgS, CuS, PbS, Bi 2 S 3 . AsjSa, Sb 2 S 3 , SnS. The last three dissolve in (NH 4 ) 2 Sz. Group 3. Precipitated by (NH 4 ) 2 S from neutral solution: FeS, CoS, NiS, MnS, ZnS, FeS + S from ferric salts, Cr(OH) 3 , A1(OH) 3 . Trivalent metals may be separated before addition of (NH 4 ) 2 S by precipitating as hydroxides with NH 4 OH in presence of NH 4 C1. Addition of (NH 4 ) 2 S to filtrate precipitates sulphides of bivalent metals. Group 4. Precipitated by (NH 4 ) 2 COa : BaCO 3 , SrCO 3 , CaCO 3 . Ammo- nium phosphate precipitates from filtrate MgNH 4 PO 4 . Solution contains K, Na, and NH 4 salts. APPENDIX 661 VII. ELECTROMOTIVE SERIES OF THE ELEMENTS Potassium Iron Silver Sodium Cobalt Palladium Barium Nickel Platinum Strontium Tin Gold Calcium Lead Magnesium Hydrogen Iodine Aluminium Copper Bromine Manganese Arsenic Oxygen (O ) Zinc Bismuth Chlorine Chromium Antimony Oxygen (OH ~) Cadmium Mercury INDEX (The references are to pages; the more important ones are given in bold- face type.) Acetic acid, 425 Acetylene, 34, 192, 423 ACHESON, 174 Acids, 41, 238, 425 action on dyes, 331 metals, 458, 460 salts, 507 organic, 425 Air, analysis, 290 content of water-vapor in, 91 liquid, 166 latent heat, 168 ALBERTUS MAGNUS, 5 Alchemy, 4 Alcohol, 422, 424 heat of vaporization, 162 wood, 425 Alkalies, 210 action on metals, 465 ALLEN-MOORE cell, 519 Alloys, 449 composition, 455 fusible, 417 Allo tropic forms of carbon, 178 of elements, 147 Alum, 571 Aluminium, 457, 466, 666 and acids, 459, 460 bronzes, 568 chloride, 570 hydroxide, 569 oxide, 569 sulphate, 570 hydrolysis, 513, 514 . tests for, 570 Alundum, 569 Alunite, 527 Amatol, 333 Ammonia, 302 absorption by charcoal, 176 action on salts, 509 chemical properties, 311 cyanamide process, 309 formation, 310 heat of formation, 306 heat of vaporization, 162 history, 303 in coal gas, 200 in water, 336 occurrence, 303 oxidation, 313, 325 physical properties, 311 preparation, 303 synthetic, 304 uses, 314 Ammonium bicarbonate, 316 carbonate, 316 chloride, 315 dichromate, 630 hydroxide, 313, 502 molybdate, 636 nitrate, 333 nitrite, 287 phosphomolybdate, 404, 636 polysulphide, 254 radical, 315 salts, 314 tests for, 316 sulphate, 316 solubility, 506 thioarsenate, 412 thioarsenite, 412 thiocyanate, 618 AMPERE, 473 Amphoteric elements, 413, 440 663 664 INDEX Aniline, 330 Antidotes for poisons, 409 Antimonic acid, 415 Antimonous acid, 415 Antimony, 413, 414 chlorides, 415 halides, 416 oxides, 415 salts, 415 sulphides, 416 Aqua regia, 329 Aqueous vapor, 85 Aragonite, 533 Argon, 295, 301 use, 298 ARRHENIUS, 215 Arsenic, 407 acids of, 409 halides, 411 pentoxide, 409 properties, 408 sulphides, 411 trioxide, 408 trisulphide, 412 Arsine, 410 Arsenite, 407 Asbestos, 432 Association of molecules, 379 Atmosphere, 284, 289 Atom, structure of, 652 Atomic numbers, 646 theory, 52 weights, determination, 345, 348, 352 Atoms, number in gram-molecule, 156 Auric. See also gold. chloride, heat of formation, 143 AVOGADRO, law of, 346 AVOGADRO'S number, 651 Azurite, 594 Babbit metal, 454, 455 BACON, ROGER, 5 Bacteria, 114, 295 nitrifying, 288 Baking powders, 524 BALARD, 365 Barium, 531, 646 chloride, 547 hydroxide, 546 solubility, 531 oxides, 546 peroxide, 150, 152 salts, 547 sulphate, 547 solubility, 531 tests for, 548 Barometer, 79 Bases, 210 Batteries, storage, 490, 493 BECQUEREL, 548 Bell metal, 455 Benzene, 199 Benzoic acid, 426 Beryllium, 551 BERZELIUS, 392 BESSEMER process, 608 BIRKELAND-EYDE process, 324 Bismuth, 417 glance, 417 ocher, 417 oxides, 418 salts, 418 sulphide, 418 tests for, 419 BLACK, 179 Blast furnace, 604 Bleaching, 113, 153, 148, 262, 280 powder, 104, 542 Blow-pipe, oxy-hydrogen, 49 Boiler scale, 554 Boiling-point, 160 of solutions, 224 Bone-black, 177 Boracite, 435 Borates, 436 Borax, 435, 436 -beads, 438, 618 hydrolysis, 514 Bordeaux mixture, 598 Boric acid, 434, 435 tests for, 439 INDEX 665 Bornite, 594 Boron, 434, 435 carbide, 435 hydrides, 439 nitride, 435 BOYLE, 80 law of, 285 BRAND, 397 Brass, 453, 455 Braunite, 639 British thermal unit, 139 Britannia metal, 455 Bromic acid, 387 Bromides, properties and test for, 369 Bromine, 365 chemical properties, 367 hydrate, 367 occurrence, 365 physical properties, 367 preparation, 366 uses, 367 Bronze, 455 Bronzes, aluminium, 568 Butyric acid, 426 By-product ovens, 199 Cadmium, 558 salts, 558 Caesium, 515 Calamine, 555 Calcite, 533 Calcium, 531 bicarbonate, 534 carbide, 191 carbonate, 533 chloride, 99, 532 hydrated, 94 cyanamide, 310 hydroxide, 538 solubility, 531 -light, 50 oxide, 537 heat of formation, 142 phosphate, solubility in acids, 507 phosphates, 540 Calcium, phosphide, 399 sulphate, 539, 554 solubility, 531 sulphide, 543 tests for, 545 Calculations, chemical, 68 Calorie, 90, 139, 159 Calorimeter, 139 Carbides, 179 Carbohydrates, 423 Carbolic acid, 330 Carbon, 171 allotropic forms, 178 burning of, temperature, 142 chemical properties, 179 cycle of, 186 physical properties, 179 dioxide, 179, 182 critical temperature, 165 heat of formation, 141 history, 179 in the air, 291 in nature, 186 occurrence, 179 physical properties, 179 preparation, 180 test for, 184 uses, 185 disulphide, 190 monoxide, 187 chemical properties, 177, 189 physical properties, 189 preparation, 188 test for, 190 prints, 630 tetrachloride, 115, 191 Carbona, 191 Carbonates, preparation, properties and test for, 184 Carborundum, 194 Carnalite, 526 CARO'S acid, 281 CASTNER-KELLNER cell, 519 Catalysis, 26, 101, 177, 307 CAVENDISH, 38, 88, 297, 323 Cell, Daniell, 479 INDEX Cell, dry, 480 electric, 479 gravity, 480 Leclanche", 481 Cellulose, 423, 596 nitrate, 427 Cement, 575 Cementite, 610 Chalcocite, 594 Chalcopyrite, 594 Chalk, 533 Charcoal, animal, 177 heat of combustion, 140 wood, 175, 176 CHARLES, law of, 82 Chemical equilibrium, 233 reactions, types of, 130 Chlorates, 384, 385 Chloric acid, 384, 385 Chlorides, behavior when heated, 115 formation, 128 test for, 128, 130 Chlorine, 99 action with water, 110 as bleaching agent, 113 chemical properties, 106 comparison of with oxygen, 115 critical temperature, 165 dioxide, 386 history, 100 hydrate, 109 liquefaction, 109 occurrence, 99 oxidizing action, 111 physical properties, 105 preparation, 100-105 test, 114 uses, 114 Chlorine monoxide, 382 Chloroform, 130 Chlorophyl, 26, 187 Chrome alum, 593, 628 Chromates, 628 Chromel, 494 Chromic anhydride, 629 chloride, 628 compounds, 627 Chromic oxide, 628 Chromium, 626 in steel, 612 test for, 635 Chromite, 626 Chromous compounds, 627 Citric acid, 426 CLAUDE, 307 Clay, 571 Climate, effect of water-vapor on, 89 Coal, 194 analysis, 195 anthracite, 196 bituminous, 196 burning of, 196 formation, 194 semi-bituminous, 196 yield of gas from, 202 Coal tar, 200 Cobalt, 619 tests for, 620 Cobaltic compounds, 619 Cobaltite, 408 Cobaltous compounds, 619 Coins, 601 Coke, 198 analysis of, 196 Colemanite, 435 Colloidal clay, 574 metals, 447 Combination, 131 Combustion, 31 heat of, 140 long-flame, 207 spontaneous, 33 surface, 206 Compounds, chemical, 17 Conduction, electrolytic, 489 Conductivity, electrical of metals, 446 heat, 446 Copper, 463, 594. See also cuprous and cupric action of acids on, 462 ferro cyanide, 614 hydroxide, 596 metallurgy, 469 nitrate, 341, 598 INDEX 667 Copper oxides, 595, 596 purification, 495 reaction with nitric acid, 327 sulphate, 93, 94, 598, 513 sulphide, 599 tests for, 599 Coral, 533 Corrosion of iron, 617 of metals, 462 Corundum, 569 Cotton, 596 COULOMB, 473 Coulometer, 483 Couples, metallic, 458, 478 COURTOIS, 370 Critical temperature, 164 Crocoisite, 627 CROOKES, 652- Croton water, analysis, 337 Cryolite, 570 Crystallography, 157 Cupric bromide, 597 chloride, 597 Cuprite, 594 Cuprous chloride, 597 cyanide, 598 iodide, 597 CURIE, 548, 652 D DALTON, 54-56, 346, 358 DAVY, 100, 338, 527 DEACON process, 101 Decomposition, double, 131, 134, 221 Deliquescence, 126 Density, 159, 442 Detonators, 331 Dew, 90 DEWAR, 109 Diamond, 140, 172, 178 Dichromates, 629 Diffusion, 155 of gases, rate, 45 Disilicic acids, 431 Distillation, 96, 163 Dissociation, 113 electrolytic, 215 Double decomposition, 131, 134, 221 DULONG, 352 DUMAS, 109 Duriron, 464, 612 Dust hi the air, 295 Dyes, bleaching, 113 sulphur, 249 E EDISON storage cell, 493 Efflorescence, 93 ELDRED, 207 Electricity, 8 Electrochemistry, 471 Electrolysis, 217 of water, 27 Electrolytes, 216 Electrolytic conduction, 489 Electromotive force of cells, 488 series of elements, 223, 486, 488 Electronic valence, 389 Electrons, 389, 472, 482, 647 Electroplating, 464, 496 Electrotypes, 496 Elements, 19 acid-forming, 439 amphoteric, 413, 440 metallic, 209 non-metallic, 210 Emery, 569 Emulsions, 437 Enamels, 437, 573 Endothermic reactions, 141 Energy, 11, 137 electrical, 472 free, 137 Epsom salt, 551 Equations, writing of, 61, 103, 132, 257-259 partial, 258 thermo-chemical, 141 Equlibrium, chemical, 233 Esters, 426 Ethane, 421 Ethyl acetate, 426 borate, 439 chloride, 338 668 INDEX Ethylene, 422 Eutectic, 451 Exothermic reactions, 141 Explosions, 47, 330, 342 Family, chemical, 413, 439, 441 FAHRENHEIT, 78 FARADAY, 109, 164, 165 law of, 482 Fats, 426 cooking, 50 heat of combustion, 140 FEHLING'S solution, 596 Feldspar, 432 Fermentation, 424, 425 Ferric acetate, 618 -ammonium alum, 511 chloride, 614 compounds, 614 hydroxide, 503, 615 oxalate, 593 oxide, 615 sulphate, 616 Ferrite, 611 Ferrous-ammonium sulphate, 613 carbonate, 613 chloride, 613 compounds, 613, 614 sulphate, 613 sulphide, 614 Fertilizers, 541 Fire extinguishers, 115, 185 Flame, Bunsen, temperature, 204 hydrogen, temperature, 142 Flames, 205 Flint, 430 Flotation, 595 Fluorides, 381 Fluorine, 375 discovery, 377 occurrence, 376 preparation, 377 properties, 378 Flux, 437, 468 Formalin, 425 Formulas, 58 Formulas and valence, 120 determination, 349 graphic, 119, 389 FOWLER'S solution, 409 Franklinite, 555 FRASCH, 246 Freezing mixtures, 505 -point of solutions, 227 Fructose, 424 Fuels, analyses, 196 heat value, 196 Fuller's earth, 574 Furnaces, electric, 494 G Galena, 582 Garnet, 431 Gas-black, 178 burning, of, 204 coal, 200 illuminating, 200 calorific value, 201 masks, 643 natural, 300, 422 producer, 202 water, 201 Gases, 155 adsorption, 176 calculation of volumes, 72 kinetic theory, 156 liquefaction, 164 measurement, 77 molecular volume, 75 reduction of volume, 84, 86 relative densities, 75 solubility, 166 Gasoline, 115, 422 GAY-LUSSAC, 322, 370, 377 law of, 345 Germanium, 578 GLADSTONE, 239 Glass, 409, 543 history, 3 and hydrofluoric acid, 381 GLAUBER, 123 GLAUBER'S salt, 124 Glucose, 424, 596 INDEX 669 Glycerine, 426 Gold, 599 tests for, 602 Gram-molecular-volume, 74 Gram-molecular-weight, 59, 74 Granite, 432 Graphite, artificial, 173, 174, 192 heat of combustion, 178 Gravity, specific, 159 Guano, 321 GULDBERG, 239 Gun metal, 455 powder, 528 Gypsum, 539 H HABER, 305, 308 HALL, 567 Halogen family, 364 and periodic law, 388 Hardness, 442 Hausmannite, 639 Heat, 7 energy, measurement, 138 of formation, 141 of chlorides, 143 of oxides, 142 of vaporization, 161 specific, 159 Helium, 299, 301, 561, 653 HELMHOLTZ, 137 Hematite, 602 Hemoglobin, 603 HENRY, law of, 166 VAN'T HOFF, law of, 242 HOFMANN apparatus, 9 Honey, 424 Horse-power, 473 Hydrates, 92, 109 Hydra zine, 317 Hydrazoic acid, 318 Hydriodic acid, 373, 374 Hydrobromic acid, 307, 368 Hydrocarbons, action with chlorine, 108 Hydrochloric acid, 123 chemical properties, 125-127 Hydrochloric acid, commercial, 127 constant boiling solution, 128 electrolysis, 100 physical properties, 124, 127 preparation, 124 Hydrocyanic acid, 529 Hydrofluoric acid, 378, 380 Hydrofluosilicic acid, 434 Hydrogen, 38 action with chlorine, 107 bromide. See hydrobromic acid. chemical properties, 46, 48 chloride. See hydrochloric acid. critical temperature, 165 explosion, 47 history, 38 in the air, 294 iodide. See hydriodic acid. molecular weight of, 156 occurrence, 38 peroxide, 149-153, 294 physical properties, 44-46 preparation, 39-44 selenide, 393 sulphide, 249, 251 test for, 254 telluride, 393 temperature of flame, 49 uses, 49 Hydrolysis, 252 of salts, 511 Hydroxides, 501 Hydroxylamine, 318 Hypo, 593 -bromous acid, 387 -chlorites, 383 -chlorous acid, 382, 383 -iodous acid, 387 -phosphorous acid, 405 Hypothesis, 57 I latrochemists, 413 Ice, heat of fusion, 167 lodic acid, 387 anhydride, 387 Iodides, test for, 375 670 INDEX Iodine, 370-372 chlorides, 387 pentoxide, 387 lodoform, 373 lonization, extent of, 228, 231 of salts, 511 Ion, common, effect of, 509 Ions, 216 Iron, 447, 603. See also ferrous and ferric. action of acids on, 460 as acid-forming element, 616 cast, 606 corrosion of, 463, 617 metallurgy, 603 passive, 460 properties, 612 purification, 495 tests for, 618 wrought, 607 Isomorphism, 503 Isotopes, 650, 656 JOULE, 473 K Kainite, 526 Kaolin, 432 Kerosene, 422 Kinetic theory, 156, 160 Krypton, 301 Lactic acid, 425 Lactose, 424 Lampblack, 178 LANDOLT, 53 LAUE, 647 LANGMUIR, 655 Lanolin, 528 LAVOISIER, 25, 52, 256, 284, 322 Law, 57 of Avogadro, 346 of Boyle, 80 of Charles, 82 of conservation of energy, 137 Law, of conservation of matter, 53 of definite proportions, 54 of Dulong and Petit, 352 of Faraday, 482 of Gay-Lussac, 345 of Henry, 166 of Le Chatelier, 243 of molecular concentration, 238 of multiple proportions, 54 of Stephan, 400 of van't Hoff, 242 periodic, 358 Lead, 582 and acids, 459 atomic weight, 649 carbonate, 585 chloride, 584 chromate, 627, 629 hydroxide, 584 iodide, 584 oxides, 583 salts, 586 storage battery, 491 tests for, 586 Leather, 571 LEBLANC, 520, 521 LE CHATELIER, 243, 306 LEWIS, G. N., 655 LIEBIG, 267 Light, 7 Lime, 207, 537 chloride of, 104 -light, 50 slaked, 538 Limestone, 533 Limonite, 603 Liquefaction of gases, 164 Liquids, 156 Lithium, 515 Lithopone, 547 LOCKYER, 299 Lodestone, 615 LULLUS, RAYMOND, 5 M Magnesia alba, 551 Magnesium, 457, 551 INDEX 671 Magnesium-ammonium phosphate, 404, 555 burning of, 400 carbonate, 554 chloride, 99, 553 hydroxide, 502, 553 oxide, 552 pyrophosphate, 555 sulphate, 553, 554 tests for, 555 Magnetite, 603 Malachite, 594, 598 Manganates, 642 Manganese, 639 dioxide, 101 in steel, 612 tests, 645 Manganic compounds, 641 Manganite, 639 Manganites, 641 Manganous compounds, 640 Marble, 533 MARSH'S test, 410 Mass law, 239 Matter, 11 MAYER, 137 MAYOW, 284 MEKER burner, 204 Melting-point, 167 MENDELEJEFP, 359 Mercuric chloride, heat of formation, 143 oxide, heat of formation, 142 Mercury, 559 cyanide, 564 fulminate, 564 halides, 561 nitrates, 560 oxides, 560 sulphides, 563 symbol for, 65 tests for, 564 Meta-acids, 264 Metallic couples, 458, 478 Metallurgy, 457, 467 Metals, action with acids, 42, 43, 222, 458 Metals, action on alkalies, 43, 464 action on salts, 222, 465 boiling- and melting-points of, 448 chemical properties, 457 colloidal, 447 corrosion, 462 electromotive series, 486 occurrence, 466 passive, 460 physical properties, 442 physical state, 446 Metastannic acid, 581 Metathesis, 131 Methane, 108, 421, 422 Methyl alcohol, 425 ether, 422 Mica, 431 Microcosmic salt, 438 Mill scale, 618 Mispickel, 408 MOISSAN, 172, 179, 377 Mol, definition of, 59 Molar solutions, 228 Molecular weight, 59 Molybdenite, 636 Molybdenum, 636 Molybdic anhydride, 636 Monel metal, 455 MORLEY, 356 Mortar, 538 MOSELEY, 647 Muriatic acid, 100 N Nascent state, 112 NELSON cell, 519 Neon, 301 Neutralization, 219 Nichrome, 494 Nickel, 465, 620 catalysis with, 50 compounds of, 620, 621 in steel, 612 tests for, 621 Nitrates, 331 in water, 336 672 INDEX Nitrates, occurrence, 321 test for, 332 Nitric acid, 321 action on skin, 331 action with metals, 327, 460 chemical properties, 326 history, 321, 322 in the air, 294 physical properties, 326 preparation, 322 synthesis, 323 test for, 332 Nitric anhydride, 343 oxide, 339 heat of formation, 323 preparation, 340 properties, 340 Nitrites, occurrence, 336 test for, 335 Nitro-benzene, 330 -cellulose, 330 -glycerine, 330 -toluene, 330 Nitrogen, 284 chemical properties, 288 history, 285 occurrence, 286 physical properties, 288 preparation, 287 uses, 289 Nitrogen chloride, 387 dioxide, 341 dioxide, chemical properties, 342 iodide, 387 oxides, 337 pentoxide, 343 tetroxide, 341 Nitrosyl chloride, 329, 341 Nitrous acid, 333, 335 anhydride, 334, 342 oxide, 337-339 Non-metals, electromotive series, 488 reactions with salts, 223 Normal solutions, 229 O Ocean, salts in, 96 Oleic acid, 426 Olein, 426 ONNES, 300 Opal, 430 Orpiment, 407, 411 Orthoclase, 432 Orthophosphates, 402-404 Orthophosphoric acid, 402 Osmium, 625 Ovens, bee-hive, 199 by-product, 199 Oxalic acid, 426 Oxidation, 49, 631, 632, 644 slow, 32 Oxides, 499 Oxygen, 21 chemical properties, 30 comparison of with chlorine, 15 critical temperature, 165 history, 20 in nature, 35 in the air, 290 preparation, 24-29 physical properties, 29 separation from air, 23 standard for atomic weights, 351 uses, 33 Ozone, 144 chemical properties, 146, 147 formation, 145-146 heat of formation, 145 history, 144 in the air, 294 physical properties, 146 preparation, 145 uses, 148 Paint-remover, 432 Palladium, 625 Palmitic acid, 426 Paper, 423 PABACELSUS, 38, 414 INDEX 673 Paraffin, 422 Paris green, 598 PASCAL, 285 PASTEUR, 285 Peat, 196 Perchlorates, 385, 386 Perchloric acid, 497 Perchloric anhydride, 386 Periodic law, 358, 388, 439 Permanganates, 643 Permanganic acid, 643 Permutit, 536 Peroxides, 149, 150 Persulphides, 280 Persulphuric acid, 280 Petroleum, 422 Pewter, 455 Phase, 169 Philosopher's stone, 4 Phlogiston, 52 Phosgene, 177 Phosphine, 405 Phosphorescence, 400 Phosphoric acids, 401 Phosphorus, 396 action with bases, 405 chemical properties, 399 discovery, 397 iodides, 407 occurrence, 397 pentachloride, 406 pentoxide, 401 physical properties, 398 preparation, 397 sulphides, 407 trichloride, 406 trioxide, 401 Phosphorous acid, 404 Photography, 282, 692 Picric acid, 330 Pitch-blende, 548 Plants, growth of, 186 Plaster of Paris, 539 Platinum, 447, 623 compounds, 624 Poisons, 409 Polymers, 265 Polysulphides, 255 Porcelain, 573 Portland cement, 575 Potentials, single, 485 Potassium, 526 carbonate, 527 chlorate, preparation, 384, 385 chloride, solubility, 506 chloroplatinate, 529 chromate, 629 cobalticyanide, 619 cyanide, 529 dichromate, 629 ferrate, 616 ferricyanide, 616 ferrite, 616 ferrocyanide, 614 hydroxide, 527 manganate, 642 metabisulphite, 264 nitrate, 528 heat of solution, 504 solubility, 505 permanganate, 642, 643 sulphate, solubility, 506 tests for, 529 Pressure, atmospheric, 79 effect on boiling-point, 168 equilibrium, 243 freezing-point, 168 electrolytic solution, 484 measurement, 78 PRIESTLEY, 21, 22, 24, 88, 179, 256, 284, 303, 314, 338 Producer-gas, 202 Propane, 421 Properties, physical, chemical, 11 variable, characteristic, 9 PROUST, 54 Prussian blue, 614 Putty, 534 Pyrolusite, 640 Pyrometers, 475 Pyrophosphoric acid, 402 Pyrosulphates, 276 674 INDEX Quartation, 599 Quartz, 430 Q R Radicals, 120, 131 Radioactivity, 646, 648 Radiometer, 8 Radium, 648, 649, 651 RAMSAY, 296, 299 RAYLEIGH, 296 Reactions, endothermic, 141 exothermic, 141 of double decomposition, 221 rate, 240 reversible, 110, 234 Realgar, 407, 411 Reduction, 48, 632, 644 Rhodium, 625 RICHARDS, 356 Rochelle salt, 596 RONTGEN, 652 ROSE'S metal, 455 Rubidium, 515 Ruby, 569 Russia iron, 464 Rusting of iron, 617 Ruthenium, 625 RUTHERFORD, 285, 652, 655 S Salt, in sea water, 96 Saltpeter, Chile, 286 Salts, 211 acid and basic, 212 complex, 510 double, 510 hydrolysis, 511 ionization, 511 properties, 503 reaction with metals, 222, 465 with non-metals, 223 solubility, 506 in acids, 507 in ammonia, 509 in other salts, 509, 510 Sandstone, 429 Sapphire, 569 SCHEELE, 100, 285 SCHONBEIN, 144 Science, 57 Selenic acid, 395 Selenium, 392 compounds, 393 dioxide, 394 Serpentine, 432 Sewage, 97 Siderite, 603 Silica, 430 Silicates, 430, 431, 466 test for, 433 Silicic acid, 430 Silicon, 428, 429 carbide, 192 dioxide, 429 hydrides, 433 tetrachloride, 433 tetrafluoride, 434 Silk, 581 artificial, 596 Silver, 588 chloride, heat of formation, 143 German, 455 halides, 591 nitrate, 591 orthoarsenate, 404 orthophosphate, 404 oxide, heat of formation, 142 reaction with hydrogen per- oxide, 153 oxides, 590 tests for, 592 Slag, 468, 606 Smalt, 619 Smaltite, 408 Smithsonite, 555 Soap, 426, 437, 535 history, 4 Soapstone, 432 SODDY, 652 Sodium, 516 amalgam, 315 arsenide, 410 INDEX 675 Sodium bicarbonate, 523 carbonate, 520, 523 hydrated, 93 chloride, 516 heat of solution, 505 in sea water, 99 dichromate, 630 fluosilicate, 525 hydroxide, 102, 618 hypochlorite, hydrolysis, 512 iodate, 387 metaborate, 437 metaphosphate, 438 nitrate, 321 solubility, 506 nitrite, 334 oxide, heat of formation, 142 permanganate, 643 peroxide, 28, 153 persulphate, 280 pyroantimonate, 416, 525 salts of, 525 silicate, 432 stearate, 426 sulphate, solubility, 506 tests for, 525 tetrathionate, 280 thiosulphate, 279 zincate, 557 Soldering, 438 Solders, 452, 453 Solids, 157 Solubility of gases in liquids, 166 of salts, 504 product, 502 Solutions, 212 boiling-point, 224 freezing-point, 227 normal, 229 SOLVAY, 520, 522 Specific gravity, 159 heat, 159 Spectroscope, 525 Spelter, 556 Sphalerite, 408, 555 Spinthariscope, 651 Stannic acid, 580 Stannic chloride, 581 Stannous chloride, 581 hydroxide, 580 Starch, 423, 525 heat of combustion, 140 as test for iodine, 114 Stearic acid, 426 Stearin, 426 Steel, 607 Steels, alloy, 611 high speed, 612 STEPHAN'S law, 400 Stereotype metal, 455 Stibine, 416 Stone, artificial, 432 Strontium, 531, 645, 646 hydroxide, 531 sulphate, solubility, 531 Stucco, 540 Sublimation, 167 Substances, pure, 15 Substitution, 108, 131 Sucrose, 424 Sugar, 424 heat of combustion, 178 Sulpharsenites. See thioarsenites. Sulphates, 277 test for, 278 Sulphides, preparation, 253 solubility in acids, 508 test for, 254 use in analysis, 508 Sulphites, 263 test for, 265 Sulphur, 245 chemical properties, 248 chlorides, 248 extraction, 246 occurrence, 245 physical properties, 247 uses, 249 Sulphur dioxide, catalytic oxidation, 270 chemical properties, 260 critical temperature, 165 history, 256 in the air, 294 676 INDEX Sulphur dioxide, occurrence, 256 physical properties, 260 preparation, 257 uses, 261, 262 Sulphur trioxide, 265 Sulphuric acid, 267 action on metals, 461 chemical properties, 275 contact process, 269 history, 267 lead chamber process, 270 manufacture, 268 physical properties, 274 uses, 279 Sulphurous acid, 263 Supercooling, 169 Superheating, 169 Sylvite, 526 Symbols, 58 Talc, 432 Tartar, cream of, 524 -emetic, 415 Telluric acid, 395 Tellurium, 392 chloride, 394 dioxide, 394 nitrate, 395 sulphate, 394 Tellurous acid, 395 Temperature, critical, 164 effect on equilibrium, 240 Tempering steel, 610 Tenacity, 444 Thermite, 568 Thermochemistry, 140 Thermometers, 77 Theory, 57 electrolytic dissociation, 215 THOMPSON, J. J., 652 Thioarsenates, 412 Thioarsenites, 412 Thiosulphates, 279 Tin, 579. See also stannous and stannic. action of nitric acid on, 461 Tin, oxides, 579 recovery, 114 sulphides, 582 symbol for, 59 test for, 582 T. N. T., 330, 333 Toluene, 199 TOWNSEND cell, 518 Transition-points, 168 Tripoli powder, 430 Tungsten, 636 in steel, 612 Type metal, 455 U Ulexite, 436 Uranium, 649 compounds, 637 Valence, 117 effect of change in, 440 effect on chemical properties, 626 electronic, 389 electrons, 653 influence on basic properties, 639 of metals and properties, 566, 578 positive and negative, 390, 632, 644 VALENTINE, BASIL, 414, 417 VAN HELMONT, 179 VAN'T-HOFF, 242 Vapor, aqueous, 85 Vaporization, heat of, 161 Vegetables, content of water, 89 Venetian red, 615 Ventilation, 148, 292 Verdigris, 598 Vinegar, 211 W WARD, 267 War-gases, 177 Water, 88 action with metals, 40 analysis, 336 INDEX 677 Water, chemical properties, 91, 92 composition, 94, 95 critical temperature, 165 dissociation by heat, 113 effect of pressure on boiling- point, 168 effect of pressure on melting- point, 168 electrolysis, 27, 39 hard, 437, 535 heat of formation, 142 from ions, 219 heat of vaporization, 162 history, 88 ionization, 512 occurrence, 88 physical properties, 89, 90 purification, 96, 177 sterilization, 148 superheated, 169 -gas, 201 vapor, 90 in the air, 292 Waters, natural, 96 WATT, 473 WELDON mud, 642 WELSBACH gas-mantels, 401 White lead, 585 Whiting, 534 WlLHELMY, 239 WOHLER, 567 Wood, 196, 425 WOOD'S metal, 417, 455 X Xenon, 301 X-Rays, 646 Zinc, 463, 465, 555 action of carbon dioxide on, 183 blende, 555 chloride, 557 hydroxide, 557 oxide, 556 salts, 557 sulphide, solubility in acids, 508 tests for, 557 THIS BOOK IS DUE ON THE LAST DATE STAMPED BELOW AN INITIAL FINE OP 25 CENTS WILL BE ASSESSED FOR FAILURE TO RETURN THIS BOOK ON THE DATE DUE. 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