COLLEGE TEXTBOOK 
 
 OF 
 
 CHEMISTRY 
 
 BY 
 
 WILLIAM A. NOYE> 
 
 DIRECTOR OF THE CHEMICAL yA 
 
 OF THE UNIVERSITY OF ILLINOIS 
 
 NEW YORK 
 HENRY HOLT AND COMPANY 
 
COPYRIGHT, 1919 
 
 BY 
 HENRY HOLT AND COMPANY 
 
 
PREFACE 
 
 This College Textbook of Chemistry is designed, more es- 
 pecially, for students of the Freshman or Sophomore years 
 in college who have not studied chemistry in the High 
 School. As with all textbooks for beginners, two purposes- 
 have been constantly kept in mind while writing the book: 
 the presentation of a few of the multitude of chemical facts 
 which touch our modern life, in such a manner that they 
 can be clearly understood, and the discussion of the theories 
 and principles around which all our chemical knowledge is 
 grouped. The teacher of chemistry is embarrassed by the 
 vast and ever increasing amount of knowledge at his dis- 
 posal and is often tempted to present many more topics than 
 the student can possibly remember. In trying to avoid 
 this difficulty many facts ordinarily included in an elemen- 
 tary textbook have been omitted and those which are given 
 are brought as far as possible into close logical relations. 
 
 The summary at the close of each chapter is a somewhat 
 unusual feature of the book. It is hoped that these summa- 
 ries will be found useful. 
 
 Success in the study of chemistry depends especially on 
 the ability to learn new facts in their relation to those which 
 have already been acquired and on the cultivation of a 
 logical as distinguished from an arbitrary memory. The 
 exercises at the close of each chapter and questions occa- 
 sionally inserted in the text are designed to assist the student 
 in this direction. These, and similar exercises prepared by 
 the teacher, should be given very careful attention in the 
 class room. 
 
 435145 
 
vi PREFACE 
 
 I wish to express my very sincere thanks to Dr. B. S. 
 Hopkins, Dr. H. G. Deming, Dr. Charles Davidson and 
 Mrs. H. A. Davidson, who have read the manuscript and 
 made many valuable and helpful criticisms; also to Dr. J. H. 
 Ready and Dr. B. S. Hopkins, who have read the proofs, 
 and to Dr. Charles Davidson, who has prepared the index. 
 
CONTENTS 
 
 CHAPTER PAGE 
 
 I. Fundamental Principles 1 
 
 II. Combustion 6 
 
 III. Hydrogen . 19 
 
 [V. Weights and Measures, Molecular Theory, Laws of 
 
 Gases 28 
 
 Composition of Water. Laws of Composition by 
 
 Weight. The Atomic Theory 42 
 
 VI. Properties and Uses of Water. Vapor Pressure ... 58 
 
 VII. Sodium, Acids, Bases, Salts 72 
 
 VIII. Hydrochloric Acid, Chlorine, Oxygen Acids of Chlorine 80 
 IX. Group VII; the Halogen Family, Bromine, Iodine, 
 
 and Fluorine 92 
 
 X. Classification of Elements. Valence 100 
 
 XI. Group VI; Sulfur, Selenium and Tellurium 109 
 
 XII. Selection of Molecular and Atomic Weights .... 130 
 
 XIII. Group V; Nitrogen 138 
 
 XIV. Air. The Noble Gases. Group Zero " 155 
 
 XV. The Periodic System 162 
 
 XVI. Group V; Phosphorus, Arsenic, Antimony and Bismuth 170 
 
 XVII. Group IV; Carbon 185 
 
 XVIII. Hydrocarbons, Gas, Flame 195 
 
 XIX. Carbon Monoxide, Carbon Dioxide, Carbon Disulfid, 
 
 Cyanides 211 
 
 XX. Carbohydrates, Alcohols, Acids, Bread, Proteins, Di- 
 gestion, Antitoxins, Alkaloids, Dyes 219 
 
 XXI. Group IV; Silicon, Tin and Lead 237 
 
 XXII. Group III; Boron, Aluminium 253 
 
 XXIII. Group II; First Division; Alkali Earth-metals, Cal- 
 cium, Strontium, Barium, Radium 260 
 
 XXIV. Metallurgy and the Preparation of Compounds of the 
 Metals 277 
 
 XXV. Group II; Second Division; Magnesium, Zinc, Cad- 
 mium and Mercury 283 
 
 XXVI. Group I; First Division; Alkali Metals: Sodium, Potas- 
 sium and Ammonium. Spectrum Analysis . . . : . 290 
 
viii CONTENTS 
 
 ^ CHAPTER PAGE 
 
 XXVII. Group I; Second Division; Copper, Silver, Gold; Pho- 
 tography 308 
 
 XXVIII. Group VI; Second Division; Chromium, Tungsten, 
 
 Uranium ...... 316 
 
 XXIX. Group VII; Second Division; Manganese 320 
 
 XXX. Group VIII; Iron, Cobalt, Nickel, Platinum .... 323 
 XXXI. Analysis .335 
 
COLLEGE TEXTBOOK OF CHEMISTRY 
 
 CHAPTER I 
 FUNDAMENTAL PRINCIPLES 
 
 Science is systematic knowledge. Especially it is that 
 sort of knowledge which discovers relations between the 
 facts or phenomena of experience and expresses these rela- 
 tions in the form of laws. 
 
 Gravitation. Laws. It is a common observation that 
 an object which is not supported will fall toward the earth. 
 Newton, one of the greatest scientific men of jthe world, 
 observed an apple fall to the ground. By reflecting on this 
 observation and by stud}dng the rates of the motions of the 
 planets about the sun he discovered the law of gravitation 
 that all material objects attract each other with a force 
 which is proportional to their masses and inversely pro- 
 portional to the squares of their distances apart. Laws 
 which have been discovered in a similar manner by ob- 
 servation, experiment and reflection furnish the nucleus 
 about which all our scientific knowledge is grouped in a 
 systematic manner. 
 
 Physical Science is that part of our systematic knowledge 
 which deals with the phenomena of inanimate nature. 
 These phenomena may be considered from two quite differ- 
 ent points of view. 
 
 (a) Physics. We may direct our attention to the motion 
 of bodies and to the phenomena of sound, heat, light and 
 electricity, or, in general, toward those phenomena which 
 
2 ; . : t . FUNDAMENTAL PRINCIPLES 
 
 we group together under the general name, energy. In its 
 simplest form energy is merely the motion of bodies or 
 some force which may produce motion. The motion of a 
 body is called kinetic energy. A force which may produce 
 motion, directly or indirectly, is called potential energy. 
 There may be some question whether there is any sort of 
 energy which is not of the nature of motion. The branch 
 of science which considers the subject of energy and its 
 transformations is called physics. 
 
 (6) Chemistry. From a quite different point of view we 
 may consider the materials of which substances are com- 
 posed, distinguishing pure, or homogeneous, substances from 
 mixtures, and studying how such substances act on each 
 other under various conditions. In other words, we may 
 direct our attention to the composition of substances. The 
 branch of science which does this is called chemistry. 
 
 Mixtures. The first step in the study of substances in a 
 chemical way is the separation of pure substances from the 
 mixtures which are met in common experience. Two of 
 the most effective means used for this purpose are crystalli- 
 zation and distillation. 
 
 The salt brines found at Syracuse, New York, in Michigan 
 and in other parts of the world contain a number of other 
 substances in solution with the salt, but when these brines 
 are evaporated very nearly pure crystals of salt separate, 
 while the other substances remain in solution. On the 
 other, hand, if the steam which is boiled away from the 
 brine is condensed, water free from salt can be obtained. 
 If a mixture of alcohol and water is boiled, the portion of the 
 mixture which distils first will contain more alcohol than 
 the portion remaining behind and by repeating the distilla- 
 tion several times it is possible to separate the original 
 mixture into nearly pure alcohol and pure water. 
 
 Pure Substances are Homogeneous, that is, all parts of 
 them are alike. If they melt or boil without decomposition 
 
LAW OF CONSTANT PROPORTION 3 
 
 each has a definite melting point or boiling point, under 
 atmospheric pressure. They always act in the same manner 
 when brought in contact with a given other substance 
 under the same conditions. If they are compounds they 
 always have the same composition by weight. 
 
 Law of Constant Proportion. The last property of pure 
 substances mentioned has been established by the accurate 
 analysis of many different compounds and these analyses 
 have established one of the most important laws of chemis- 
 try the law of constant proportion. 1 
 
 Compounds. Synthesis. If fine iron filings and sulfur 
 are thoroughly mixed in a mortar a mixture is obtained 
 which has a uniform gray color and which appears to the eye 
 homogeneous. With a microscope, however, it is easy to 
 see that the substance is still a mixture of particles of iron 
 with particles of sulfur. With a magnet the particles 
 of iron may be picked up and the particles of sulfur which 
 still cling to the iron may be blown away. If the mixture 
 is heated over a flame, after a short time it will begin to 
 glow and the mass will become red hot. After cooling, the 
 material is homogeneous and neither particles of iron nor 
 particles of sulfur can be seen with a microscope. A mag- 
 net will no longer attract the iron which the substance 
 contains. The new substance is evidently formed by the 
 union of iron and sulfur. It is called iron sulfide and is 
 spoken of by chemists as a compound of iron and sulfur. 
 Such a process of putting two substances together and 
 making from them a different, new substance is called 
 synthesis. 
 
 Direct Analysis. If oxide of mercury is heated in a test- 
 tube it will be decomposed into metallic mercury, which 
 will condense on the walls of the tube, and a gas, oxygen, 
 which will cause a glowing splinter to burst into flame. 
 By this process the compound, oxide of mercury, is decom- 
 
 1 Compounds of isotopes are an exception to this law, see p. 273.- 
 
4 FUNDAMENTAL PRINCIPLES 
 
 posed into two new substances, mercury and oxygen, and 
 the process may be called a direct analysis of the oxide of 
 mercury. The object of any chemical analysis is, of course, 
 to determine the substances which enter into the com- 
 position of some compound or mixture. 
 
 Elements. Compounds. Nearly all of the substances 
 which we find in the world, or which may be prepared by 
 artificial means more than one hundred thousand are 
 known to exist may be separated into two or more different 
 substances, just as oxide of mercury may be separated into 
 mercury and oxygen, though in most cases the methods 
 required for their separation are not so simple. Many 
 substances may be prepared by synthesis as sulfide of iron 
 is prepared from sulfur and iron. But a comparatively 
 small number of substances have been found which no one 
 has been able to separate into two or more other substances 
 and which no one has been able to prepare by putting two 
 or more other substances together. Such substances are 
 called elements. 
 
 The elements which have been longest known and which 
 are most familiar in our daily experience are the metals 
 iron, copper, tin, zinc, lead, mercury, silver and gold. 
 Another set of elements, which were not known till compara- 
 tively modern times, includes gases, oxygen and nitrogen 
 of the air, hydrogen, obtained by decomposing water, and 
 chlorine, a constituent of salt. Sulfur, which is used in 
 making sulfuric acid, or oil of vitriol, phosphorus, used in 
 matches, and carbon, which is nearly pure in charcoal, 
 graphite and diamonds, are also known to most persons 
 before they begin the systematic study of chemistry. If we 
 add to these silicon, found in sand, aluminium, an element 
 found in clay, calcium, found in limestones, magnesium, 
 also found in many limestones, sodium from salt, and potas- 
 sium from wood ashes, we have a list of twenty-one elements. 
 About ninety eight per cent of the part of the earth which 
 
CHEMICAL ELEMENTS 5 
 
 we can examine and of the air above is composed of these 
 elements, and very few others require more than a passing 
 mention in an elementary textbook. 
 
 About sixty other elements have been discovered and 
 we now have strong reasons for believing that the number 
 -of elements in the universe does not much exceed one 
 hundred. 
 
 SUMMARY 
 
 Science is systematic knowledge. 
 
 A law is a general relation between facts or phenomena. 
 
 Physics deals primarily with energy and its transfer. 
 
 Chemistry considers the composition of substances. 
 
 Pure substances are most often obtained by crystalliza- 
 tion or distillation. 
 
 Pure substances are homogeneous, have a constant melt- 
 ing point or boiling point, always act in the same way under 
 the same conditions, and have a constant composition. 
 
 Compounds are substances which can be separated into 
 two or more other substances. 
 
 Elements cannot be so separated. 
 
 The composition of a compound may be determined by 
 synthesis, by direct analysis,' or by indirect analysis. 
 
 Fewer than one hundred elements are known and of these 
 about twenty are of first importance. 
 
 QUESTIONS 
 
 1. What other substances beside salt are purified by crystal- 
 lization? 
 
 2. What substances beside water and alcohol are separated by 
 distillation? 
 
 3. What elements are contained in the following common 
 substances: starch, coal, brass, solder, kerosene, albumin, and, 
 snow, air? 
 
 4. Which is the most abundant element? Which is second? 
 Which is most important for life? 
 
CHAPTER II 
 COMBUSTION 
 
 Fire. It is a common experience to see wood or coal 
 burn, leaving ashes that are very much lighter than the 
 material which was burned. The oil of lamps and the 
 paraffin or tallow of candles seem to disappear as they 
 burn, leaving no visible products. These phenomena have 
 been familiar during the hundreds of thousands of years 
 which have passed since the human race first learned the 
 use of fire, but it was less than a century and a half ago that 
 a satisfactory explanation of burning, which we call combus- 
 tion, was discovered so very recent, in comparison with 
 the age of the race, is our knowledge of some of the most 
 fundamental facts of chemistry. 
 
 Relation of Air to Combustion. It was finally shown that 
 the disappearance of the wood or oil is not due to the de- 
 struction of the matter which they contain, but that, instead 
 of this, invisible substances are formed which weigh more 
 than those burned. Most people have noticed the film of 
 moisture which is formed on a cold lamp chimney when the 
 lamp is first lighted. This is due to water which is formed 
 by the burning oil. By drawing the air from above a burn- 
 ing lamp through clear lime water it can easily be shown 
 that another substance, carbon dioxide, is also formed. 
 If these two substances are absorbed by soda lime and 
 weighed, as may be done with the apparatus shown in 
 the figure (Fig. 1), it will be found that they weigh more 
 than the oil or the part of the candle which has been burned. 
 A further study of the matter has shown that the burning 
 
 6 
 
LAVOISIER'S EXPERIMENT 7 
 
 oil or candle takes up oxygen from the air and that if all 
 of the water and carbon dioxide are collected they weigh 
 exactly as much more than the oil or candle as the weight 
 of the oxygen which has been taken up. The substances 
 which disappear in burning change their form but there is 
 no change in their weight. 
 
 FIG. 1. 
 
 Lavoisier's Experiment. In 1775 the French chemist, 
 Lavoisier, showed clearly for the first time the relation 
 between air and combustion. Shortly before that Priestly, 
 an Englishman, had prepared oxygen gas by heating a red 
 compound, which we now call oxide of mercury. He had 
 shown that substances which burn in air burn with much 
 greater brilliancy in oyxgen gas. Lavoisier, after hearing 
 from Priestly of this experiment, devised the apparatus 
 shown in the figure. He placed some mercury in the retort, 
 A, and bent the wide neck of the retort in such a manner 
 that the bell-jar, FG, might have its rim below the surface 
 of mercury in the dish, S, while air could still pass freely 
 
8 
 
 COMBUSTION 
 
 from the interior of the retort to the interior of the bell-jar. 
 He then heated the mercury in the retort by means of the 
 charcoal furnace, MN, for several weeks. Red oxide of 
 mercury gradually accumulated on the surface of the 
 mercury in the retort, while the mercury rose in the bell-jar, 
 showing that a part of the air was disappearing. After 
 some weeks the mercury rose no further on continued 
 heating and the experiment was stopped. Lavoisier then 
 noticed carefully how far the mercury had risen and how 
 much of the air had been absorbed. When the oxide of 
 
 M 
 
 FIG. 2. 
 
 mercury, which he collected^carefully, was heated till it was 
 all decomposed, he found that he had obtained a volume of 
 oxygen gas equal to the decrease in volume in the air of the 
 retort and bell-jar. In this way it was shown that about 
 one-fifth of the volume of the air consists of oxygen and 
 four-fifths of some other gas. which does not act on the 
 mercury or support combustion. This other gas is chiefly 
 nitrogen. 
 
 Preparation of Oxygen. The preparation of oxygen by 
 heating oxide of mercury has been referred to. This method 
 of getting the gas is tedious and expensive and it is only 
 
CATALYSIS 
 
 9 
 
 used on a small scale in the laboratory or lecture room, 
 because of its simplicity and because of the historical 
 interest. For ordinary laboratory uses, small amounts 
 of oxygen are prepared by heating potassium chlorate, 
 a white, crystalline salt, which melts rather easily when 
 heated. It decomposes into potassium chloride and oxygen, 
 
 slowly at its melting point, more rapidly at ^ 
 
 higher temperatures. If some manganese 
 dioxide is mixed with the potassium chlorate, 
 the decomposition is very much hastened, 
 though the manganese dioxide is not, itself, 
 permanently changed. To prepare oxygen 
 in the laboratory a mixture of potassium 
 chlorate and manganese dioxide is heated 
 in a test-tube, or in a glass retort or flask, 
 or in a copper flask. 
 
 Catalysis. Manganese dioxide hastens 
 the decomposition of potassium chlorate 
 but is not permanently changed in the 
 process. The action can be easily shown 
 by melting some potassium chlorate in a 
 test-tube and dropping in some of the pow- 
 dered dioxide. A vigorous decomposition 
 of the chlorate will set in at once. A 
 substance which hastens a chemical action by its presence, 
 in this manner, is called a catalytic agent or a catalyst. 
 
 Collection and Storage of Gases. Gases like oxygen, 
 which are only slightly soluble in water, are collected in 
 bottles or in a gasometer (Fig. 3). The gasometer is filled 
 with water by opening the stopcocks A, allowing the water 
 to run to the bottom, and B, which allows the air in the 
 gasometer to escape. When the gasometer is full of water 
 the two stopcocks are closed, the cap C is removed and the 
 tube which delivers the oxygen is inserted. As the opening 
 closed by the cap slants upward, air cannot enter but 
 
 FIG. 3. 
 
10 COMBUSTION 
 
 the water escapes through it as the water in the upper part 
 is displaced by the gas. After filling, the cap is replaced 
 and the stopcock A is opened. The gas may then be drawn 
 off through B. 
 
 Oxygen from Liquid Air. As shown by Lavoisier's 
 experiment, air is composed mainly of oxygen and nitrogen, 
 about one-fifth of its volume being oxygen. When air 
 which has been compressed is allowed to expand, it grows 
 cold. Linde and others have devised machines by means of 
 which air can be compressed to a pressure of 2000 to 3000 
 pounds to the square inch. The air is then allowed to 
 expand through a small copper tube several hundred feet 
 in length. This is so arranged that the cold air escaping 
 from the lower end passes back over the outside of the 
 tube and thus cools still more the current of air passing 
 down the tube. In this way the air soon becomes so cold 
 that part of it becomes liquid. When this liquid air, 
 which is at a temperature of about 185 C. (or 300 F.) below 
 zero, is allowed to boil, the nitrogen, whose boiling point 
 is lower than that of oxygen, boils away first and a residue of 
 nearly pure oxygen is finally left. This method is now used 
 on a large scale as the cheapest method of preparing oxygen 
 for medicinal and commercial uses. 
 
 Properties of Oxygen. Oxygen is a colorless and odorless 
 gas. It is slightly heavier than air and one-seventh heavier 
 than nitrogen, the gas which forms about four-fifths of 
 the volume of air. 
 
 The most striking property of oxygen is its effect on 
 burning substances. A splinter of wood with a live coal 
 at the end will burst into flame in the gas. 
 
 A piece of charcoal which has been ignited will glow 
 brilliantly, and will be surrounded with a small blue flame, 
 scarcely visible at a short distance. As the charcoal dis- 
 appears, an invisible gas, carbon dioxide, takes the place 
 of the oxygen. The presence of the gas can be shown by 
 
BURNING IN OXYGEN 
 
 11 
 
 pouring some clear lime water into the -bottle in which the 
 charcoal was burned. The lime water will become turbid. 
 The precipitate formed is calcium carbonate, the same 
 substance which is found in marble and limestone. Lime 
 is manufactured by heating limestone in a kiln till the 
 carbon dioxide which it contains is expelled. The lime is 
 "slaked" by mixing it with water, and lime water is a solu- 
 tion of this slaked lime in water. The carbon dioxide of 
 the candle finally converts the lime back to the same com- 
 pound from which the lime was made in the lime kiln. 
 
 Sulfur burns in oxygen with a beautiful blue flame and the 
 oxygen is replaced by a colorless gas having a suffocating 
 odor. This gas is sulfur dioxide. It dis- 
 solves in water, imparting to it an acid 
 taste, and the solution will change the color 
 of a solution of litmus from blue to red. 
 Other acids, such as the acid in vinegar, in 
 a lemon or in cream of tartar, will change 
 the color of litmus in the same way. All 
 of these acids have a sour taste. 
 
 Phosphorus burns in oxygen with a very 
 brilliant white light. If the oxygen is dry, 
 a white powder settles on the bottom and 
 sides of the bottle in which the phosphorus 
 is burned. This powder is phosphorus 
 pentoxide. It has a very strong affinity 
 for water and will absorb water from ordinary air, com- 
 bining with the water and finally dissolving in it, giving the 
 solution an acid taste. The solution reddens blue litmus. 
 
 An iron wire or a coiled steel watch spring, which is 
 ignited by means of a little sulfur or by a string which has 
 been dipped in paraffin or sulfur, will burn in oxygen 
 (Fig. 4), throwing off sparks and forming white-hot, molten 
 globules of the magnetic oxide of iron, which drop from 
 the end. This oxide does not dissolve in water. 
 
 m 
 
 FIG. 4. 
 
12 COMBUSTION 
 
 Weight of Products Formed. Indestructibility of Matter. 
 
 - It is quite easy to show that the carbon dioxide, the sulfur 
 dioxide, the phosphorus pentoxide and the oxide of iron 
 each weighs more than the charcoal, sulfur, phosphorus or 
 iron from which it is formed. By weighing the charcoal or 
 other substance burned, the oxygen which is used up in the 
 process of burning and the carbon dioxide or other com- 
 pound which is formed, it has been shown that the weight of 
 the product formed by the combustion is always exactly 
 equal to the weight of the substance burned plus the weight 
 of the oxygen with which it combines. An examination 
 of very many different kinds of chemical action has shown 
 that, no matter how much the substances which interact 
 with each other may be altered in their appearance and 
 properties, it has never been possible to discover any 
 change in the weight of the materials involved, provided 
 that we consider the weights of all the substances which 
 enter into the action and all of the products formed. This 
 is known as the law of the indestructibility of matter, and it is 
 one of the most important and fundamental of the physical 
 laws. It is a law which is always assumed in the quantita- 
 tive study of any chemical problem. 
 
 Occurrence of Oxygen. Oxygen is the most abundant of 
 the elements and forms about one-half of that part of the 
 earth which we can examine. As shown by Lavoisier's 
 experiment, it is found free or uncombined in the air, 
 forming about one-fifth of its volume. It forms eight- 
 ninths of the weight of water and nearly one-half of the 
 weight of limestone and more than one-half of the weight of 
 sand. It enters into the composition of a large part of the 
 compounds which are known. 
 
 Kindling Temperature. Different substances vary 
 greatly in regard to the temperature at which they will 
 take fire and begin to burn rapidly. Phosphorus takes fire 
 .at a very low temperature and this property is utilized in 
 
SLOW OXIDATION 13 
 
 matches, which may be ignited by the heat produced 
 by gentle friction. The phosphorus is used to ignite, in 
 turn, sulfur or some other easily combustible substance 
 and the latter in burning generates enough heat to raise 
 the wood to its kindling temperature. Somewhat similar 
 phenomena are familiar to everyone in starting a fire. 
 
 Heat of Combustion. Various forms of apparatus have 
 been devised in which it is possible to burn charcoal, sulfur, 
 coal, meaty bread and other substances under conditions 
 such that the heat produced by their combustion can be 
 absorbed by water and the quantity of the heat generated 
 can be accurately measured. When this is done it is found 
 that the heat produced by burning a kilogram of charcoal 
 is nearly four times as great as that produced by burning a 
 kilogram of sulfur, and about one-third greater than that 
 given out in burning a kilogram of phosphorus. A low 
 kindling temperature seems to indicate a strong affinity 
 for oxygen, but a substance with a high kindling temperature 
 may give more heat in burning. 
 
 Slow Oxidation. Meat, bread, and other articles of 
 food may be burned in the apparatus which has been re- 
 ferred to, and the heat of combustion determined. In 
 this rapid burning the food is converted chiefly into carbon 
 dioxide and water. If these same articles of food are eaten 
 the processes which go on in our bodies also convert them 
 into carbon dioxide and water, the oxygen for the purpose 
 being taken in as we breathe. As the temperature of our 
 bodies is usually higher than that of the surrounding air, heat 
 must constantly escape from the surface of our bodies. 
 This suggests that the heat necessary to maintain body 
 temperature is derived from the slow oxidation of the food 
 which we eat. By means of a complicated apparatus 
 called a respiration calorimeter the heat given off from a 
 man's body during several days has been carefully measured 
 and compared with the he.at produced by burning food of 
 
14 COMBUSTION 
 
 the same kind as that which he had eaten. In this way it 
 has been demonstrated that food which is eaten and then 
 oxidized slowly in the body gives just the same quantity of 
 heat as the same food when burned rapidly. 
 
 Vegetable and animal substances which are exposed to the 
 action of bacteria in the soil or elsewhere are rapidly de- 
 stroyed by processes of oxidation somewhat similar to those 
 which go on in our bodies. 
 
 When iron or steel is exposed to the action of water and 
 air, it becomes covered with rust by a process of slow oxida- 
 tion. The rust contains hydrogen as well as oxygen and is 
 quite different in composition from the magnetic oxide of 
 iron formed by the combustion of iron wire in oxygen gas. 
 The oxide of iron formed on the surface of red hot iron ex- 
 posed to the air is, 'however, the same as that formed by 
 burning iron. 
 
 Spontaneous Combustion. A mass of oily cotton waste, 
 such as is used in cleaning machinery, or a pile of moist 
 coal, will often oxidize slowly in the air at ordinary tempera- 
 tures. The heat generated in this manner may cause the 
 temperature to rise till the kindling temperature of the 
 material is reached, when the mass will burst into flame. 
 This process is known as spontaneous combustion. The 
 spontaneous combustion of coal may be avoided either by 
 keeping it dry or by completely covering it with water. 
 
 Conservation of Energy. It has been pointed out that 
 the oxidation of the food which we eat furnishes the heat 
 to maintain the temperature of our bodies. It also fur- 
 nishes the muscular energy with which we move. The 
 chemical energy of the oxygen and of the elements or sub- 
 stances which combine with it is transformed into heat 
 energy and muscular energy. As is well known, a part of 
 the chemical energy of oxygen and of coal, oil or gas may 
 be converted into mechanical energy by a steam engine, 
 and the mechanical energy of the engine may be converted 
 
NOMENCLATURE. OZONE 
 
 15 
 
 to electrical energy by a dynamo or back to heat by means 
 of friction. The relations between these different forms 
 of energy have been very carefully measured and it has 
 been demonstrated that whenever one kind of energy 
 disappears an exactly equivalent amount of some other 
 form of energy takes its place. Energy cannot be created 
 or destroyed by any process known to us. This is the law 
 of the conservation of energy. It might be called the law 
 of the indestructibility of energy. 
 
 Names of Binary Compounds. It will have been noticed 
 that the compounds formed by the union of oxygen with 
 other elements are called oxides. Other compounds of two 
 elements are given names with the same ending, ide, 
 added to the name of the non-metallic element. . Com- 
 pounds of sulfur are called sulfides, compounds of chlorine, 
 chlorides, compounds of bromine, bromides. The prefixes 
 in such names as dioxide and pentoxide refer to the amount 
 of oxygen in the compound and will be explained later. 
 
 FIG. 5. 
 
 Ozone. If the oxygen is subjected to t^e action of a 
 silent electric discharge in the apparatus shown in Fig. 5, 
 it acquires a peculiar odor and becomes very much more 
 active, combining readily with silver and with other sub- 
 stances which are not affected by ordinary oxygen. If 
 this changed oxygen is liquefied by cooling it to a very 
 low temperature, it is found to be of a dark blue color, and 
 
16 COMBUSTION 
 
 on allowing the liquefied gas to boil away the last portions 
 of the liquid will boil at a higher temperature than ordinary 
 oxygen and will give a gas which is one-half, heavier than 
 oxygen. If this heavy gas is heated, however, it is changed 
 back completely to oxygen and no other substance can be 
 found in the product. 
 
 This active, heavy form of oxygen is called " ozone. " 
 It is a powerful germicide and is sometimes used to sterilize 
 water which has been contaminated with disease germs. 
 It is possible that the ozone formed in thunder storms 
 has some effect in purifying the air, but it is rather doubtful 
 if this is of much practical importance. 
 
 Allotropic Forms of Elements. In defining elements and 
 compounds it was implied that a compound may always 
 be separated into two or more substances which we call 
 elements and that these elements cannot be mutually converted, 
 the one into the other. Oxygen may be changed into a mix- 
 ture of ozone and oxygen, but we still call it an element 
 because the ozone may be completely changed back into 
 oxygen and there is no change of weight accompanying 
 the process. Several other elements exist in two or more 
 different forms. These are called allotropic forms. Thus 
 ozone is an allotropic form of oxygen. Ordinary phos- 
 phorus and red phosphorus are allotropic forms of that 
 element. 
 
 SUMMARY 
 
 Substances are changed in form but not destroyed by 
 burning. 
 
 The products formed by burning wood, coal or other 
 substances weigh more than the substances burned. 
 
 Lavoisier showed that mercury, when heated in the air, 
 takes up oxygen and that oxygen forms about one-fifth of 
 the volume of the air. 
 
 Oxygen may be prepared by heating oxide of mercury, 
 
SUMMARY. OXYGEN I/ 
 
 or by heating a mixture of potassium chlorate and manga- 
 nese dioxide ; also by the fractional distillation of liquid air. 
 
 Carbon, sulfur, phosphorus and iron burn in oxygen, 
 giving carbon dioxide, sulfur dioxide, phosphorus pentoxide 
 and magnetic oxide of iron. 
 
 Carbon dioxide gives a precipitate of calcium carbonate 
 with lime water. 
 
 Sulfur dioxide and phosphorus pentoxide give acids with 
 water. 
 
 The weight of the substance burned added to the weight 
 of the oxygen with which it combines is exactly equal to 
 the weight of the compound formed. Matter is neither 
 destroyed nor created by any chemical process. 
 
 Compounds can be formed from elements or separated 
 into elements, and always weigh more than any single ele- 
 ment which they contain. 
 
 Oxygen is the most abundant element known in com- 
 pounds on the surface of the earth. 
 
 Kindling temperature is the temperature at which a 
 substance takes fire and burns rapidly. 
 
 Heat of combustion is the heat generated by burning a 
 substance. It is the same whether the substance burns 
 rapidly or oxidizes slowly. 
 
 Slow oxidation may generate enough heat to cause spon- 
 taneous combustion. 
 
 Energy can neither be created nor destroyed. 
 
 Elements may sometimes be changed to allotropic forms, 
 but this always occurs without change in weight. 
 
 Ozone is an active, allotropic form of oxygen. 
 
 EXERCISES 
 
 1. Would the kindling temperature be the same in oxygen as 
 in air? 
 
 2. About 39 per cent of the weight of potassium chlorate is 
 oxygen. One liter of oxygen weighs 1.429 grams. How many 
 
 2 
 
18 COMBUSTION 
 
 grams of potassium chlorate will be required to give one liter of 
 oxygen? 
 
 3. Thirty-two parts of oxygen are required to burn 12 parts of 
 carbon. How many grams will be required to burn one pound 
 (453 grams) of pure charcoal? 
 
 4. Eighty grams of oxygen are required to burn 62 grams of 
 phosphorus. How many grams will be required to burn one 
 pound? 
 
 6. Assuming that one-fifth of the volume of the air is oxygen, 
 how many liters of air will be required to burn a pound of charcoal? 
 How many liters will be required to burn a pound of phosphorus? 
 
 6. One part of sulfur requires one part of oxygen for its com- 
 bustion and 168 grams of iron require 64 grams of oxygen. How 
 manj'' grams of oxygen and how many liters of air are required to 
 burn a pound of each? 
 
 7. What is the "oxone" method of preparing oxygen? 
 
 8. How may oxygen be prepared by means of barium peroxide 
 (Erin's Process)? 
 
CHAPTER III 
 HYDROGEN 
 
 Decomposition of Water by Iron. If iron is placed in an 
 iron or glass tube as shown in the figure, and steam is passed 
 over it while it is heated red hot, the iron will increase in 
 weight and will be gradually converted into magnetic 
 oxide of iron having the same composition as the oxide of 
 iron formed when an iron wire is burned in oxygen. At the 
 
 FIG. 6. 
 
 same time a very light gas, hydrogen, will collect in the 
 cylinder filled with water, which is placed over the tube 
 through which the excess of steam escapes. The experi- 
 ment demonstrates that water contains the elements oxygen 
 and hydrogen. It can be easily shown that water is formed 
 by burning hydrogen in oxygen and this completes the 
 proof that water is a compound of oxygen and hydrogen 
 and of nothing else. 
 
 Decomposition of Water by Sodium or Potassium. 
 A piece of sodium wrapped in paper and thrust quickly 
 
 19 
 
20 HYDROGEN 
 
 under the mouth of an inverted cylinder filled with water 
 will act rapidly on the water, liberating hydrogen. If the 
 sodium is thrown on the surface of the water, the same will 
 occur, but the hydrogen will escape. If the sodium is put 
 on a piece of paper floating on the surface of the water, to 
 prevent its rolling around, enough heat will be generated 
 so that the liberated hydrogen will take fire and burn. 
 The product formed by the action of the sodium on the 
 water dissolves in the water. When the solution is evapo- 
 rated a white solid remains, which is composed of sodium, 
 hydrogen and oxygen. This is called, 
 because of its composition, sodium 
 hydroxide. The name is made up from 
 the names of the three elements which 
 it contains. The presence of the sodium 
 hydroxide can also be shown without 
 evaporation, by the action on red litmus 
 paper, which is turned blue by the 
 solution. 
 
 Potassium acts on water, in a similar 
 manner but so much heat is generated 
 that the hydrogen catches fire without 
 the use of the paper. The yellow color of the flame of the 
 hydrogen from the sodium and the violet color of that from 
 the, potassium are due to the metals and not to the hydrogen. 
 Pure hydrogen burns with an almost invisible flame. 
 
 We are familiar with many articles of food which have a 
 sharp, acid taste. The most familiar of these are vinegar 
 which contains acetic acid, and lemons which contain citric 
 acid. It has been pointed out (p. 11) that the solutions 
 formed by dissolving sulfur dioxide and phosphorus pentox- 
 ide in water also have an acid taste. In both cases the 
 taste is not due to the oxide, but to a compound of the oxide 
 with water. Many other substances give solutions with 
 this same characteristic, acid taste and it has been found 
 
 FIG. 7. 
 
PREPARATION OF HYDROGEN 
 
 21 
 
 
 that all such substances contain hydrogen. Two very common 
 and important compounds of this sort are hydrochloric acid, 
 a compound of hydrogen and chlorine, and sulfuric acid, 
 a compound of hydrogen with sulfur and oxygen. 
 
 Hydrogen from Acids and Metals. If a dilute solution 
 of these acids is brought into contact with zinc or iron 
 the hydrogen of the acid is displaced by the metal and 
 liberated, very much as the hydrogen of 
 water is displaced by sodium or potassium. 
 If hydrochloric acid and zinc are used, the 
 zinc combines with the chlorine to form 
 zinc chloride. If sulfuric acid is used, the 
 zinc combines with the sulfur and oxygen, 
 giving zinc sulfate. When iron is used, the 
 compounds formed are ferrous chloride and 
 ferrous sulfate. 
 
 Substances such as zinc chloride and 
 zinc sulfarte, ferrous chloride and ferrous 
 sulfate, formed by replacing the hydrogen 
 of an acid by a metal, are called salts. 
 
 The preparation of hydrogen from zinc and dilute hydro- 
 chloric or sulfuric acid may be carried out by means of the 
 simple apparatus shown in Fig. 8. Larger quantities of 
 hydrogen may be obtained conveniently by means of a 
 Kipp generator (Fig. 9). The zinc is placed in the middle 
 bulb and the dilute acid is poured in through the upper 
 bulb, which communicates with the lower one through the 
 tube A. When the stopcock B is opened, the acid rises and 
 comes in contact with the zinc in the middle bulb and the 
 generation of hydrogen begins. Whenever the stopcock 
 is closed the pressure of the hydrogen generated forces the 
 acid away from the zinc and the action ceases as soon as the 
 acid moistening the surface of the zinc is exhausted. 
 
 In using any hydrogen generator, because of the explosive 
 character of mixtures of hydrogen and air, it is necessary 
 
 FIG. 8. 
 
22 
 
 HYDROGEN 
 
 to be careful not to bring a flame near the exit tube before 
 the air in the generator is displaced by hydrogen. Whether 
 the air has been expelled sufficiently for safety can be 
 determined by filling a test-tube with the gas in the manner 
 shown in Fig. 10. After filling the tube and removing it to 
 a safe distance from the exit tube, with the mouth of the 
 test-tube always held downward, a lighted match or flame 
 may be applied to its mouth. If there is a rather sharp 
 
 FIG. 9. 
 
 FIG. 10. 
 
 explosion, extending up into the tube, there is still enough 
 air mixed with the hydrogen to render it dangerous. If 
 there is only a slight explosion and the hydrogen continues 
 to burn in the tube at the surface where the gas is in contact 
 with the air, the hydrogen is pure enough to be safe. 
 
 Properties of Hydrogen. Hydrogen is the lightest gas 
 known. One liter at the freezing point of water (0), 
 and under normal atmospheric pressure, equal to that of 
 760 mm. of mercury, weighs 0.09 gram, while a liter of air 
 weighs 1.293 grams, about 14^ times .as much. A liter 
 
REDUCTION 23 
 
 of oxygen weighs 1.429 grams, nearly 16 times as much 
 as a liter of hydrogen. Because of its lightness hydrogen 
 is used to fill balloons. Soap-bubbles filled with the gas 
 rise through the air. 
 
 Hydrogen may be liquefied by a process similar to that 
 which has been described for liquefying air. The com- 
 pressed gas must, however, be cooled with liquid air before 
 it is allowed to expand through the spiral tube. Liquid 
 hydrogen is very light indeed and boils at a temperature 
 about 65 below the boiling point of liquid air. By causing 
 the liquid hydrogen to boil under diminished pressure a part 
 of it may be frozen to a solid, at 252.5. 
 
 The most striking chemical property of hydrogen is its 
 strong affinity for oxygen. A jet of hydrogen will burn 
 quietly in air or in oxygen and moisture will be deposited 
 on a cold surface held over the flame. By collecting the 
 water formed by burning dry hydrogen in dry air and 
 determining its freezing point and boiling point, it can be 
 shown that it is really pure water, and this proves that water 
 is composed of two elements hydrogen and oxygen. 
 
 Mixtures of hydrogen and air explode when ignited and 
 mixtures of hydrogen and oxygen explode still more 
 violently. 
 
 Occurrence of Hydrogen. The occurrence of hydrogen 
 in water and in acids has been mentioned. It is also a con- 
 stituent of practically all organic compounds, that is, of the 
 compounds of carbon which are found in vegetable and 
 animal bodies, in coal, petroleum, natural gas and in many 
 thousands of carbon compounds which are made in factories 
 for use as dyes and medicines. There is a very minute 
 quantity of free hydrogen in the air, and it is believed 
 that the outermost parts of the atmosphere are nearly 
 pure hydrogen. 
 
 Reduction. Hydrogen will not only burn in air but it will 
 also take oxygen away from many oxides when these are 
 
24 HYDROGEN 
 
 heated in a current of the gas. Copper oxide when heated 
 gives up its oxygen very readily to hydrogen which is passed 
 over it, and is changed to metallic copper. Such a removal 
 of oxygen is called reduction and the oxide of copper is 
 said to be reduced to metallic copper. At the same time 
 the hydrogen is oxidized to water. It is evident that the 
 oxidation of the hydrogen and the reduction of the copper 
 oxide are opposite processes. In one oxygen is added; 
 in the other it is removed. In almost all cases when one 
 substance is reduced some reducing agent is at the same 
 time oxidized. 
 
 Reversible Reactions. If hydrogen is passed over heated 
 magnetic oxide of iron a part of the hydrogen will be 
 oxidized to water and the oxide of iron will be slowly re- 
 duced to metallic iron. This seems rather surprising 
 because it has been stated before that when steam is 
 passed over heated iron the iron is oxidized and hydrogen is 
 liberated. It is natural, at first, to think that the oxidation 
 of the iron by the steam takes place because the affinity of 
 the iron for oxygen is greater than that of hydrogen. 
 When we discover, however, that hydrogen can take oxygen 
 away from oxide of iron it becomes evident that this 
 simple view of chemical affinity cannot be true. Since 
 the iron is oxidized when the steam is in excess and the 
 hydrogen is constantly removed from the point of action, 
 while the oxide is reduced when the hydrogen is in excess 
 and the steam is all of the time carried away by the current 
 of the gas, it must be that the reaction is reversible. As 
 long as steam, hydrogen, iron and oxide of iron are all present 
 the action is all of the time going in both directions, but 
 if the steam is in excess the action goes faster in the direction 
 toward the formation of oxide of iron and hydrogen and 
 it may finally become complete in that direction. If the 
 hydrogen is in excess, it goes faster in the other direction 
 and the oxide may be completely reduced to metallic iron. 
 
OXYHYDROGEN BLOWPIPE 25 
 
 Oxyhydrogen Blowpipe. This is an instrument by 
 means of which oxygen and hydrogen may be brought 
 together and burned, giving an intensely hot flame. The 
 temperature may be as high as 2000 to 2500 . 1 The tem- 
 perature of the flame is limited by the fact that the 
 combination of the oxygen and hydrogen is a reversible 
 reaction. At moderate temperatures and even at a tem- 
 perature of 2000 the reaction goes almost exclusively in 
 the direction toward the formation of water. At very high 
 
 FIG. 11. 
 
 temperatures water is partly decomposed into oxygen and 
 hydrogen and at such temperatures there will be a mixture 
 of oxygen, hydrogen and water. As the temperature 
 rises the proportion of water in the mixture will decrease, 
 and, since the heat comes from the union of the oxygen and 
 hydrogen, the temperature which can be obtained by burn- 
 ing the mixed gases is limited because at very high tem- 
 peratures a considerable part of the oxygen and hydrogen 
 remain uncombined. At 3700 a mixture of oxygen, 
 hydrogen and water will contain only 60 per cent of its 
 weight as water, 40 per cent of the oxygen and hydrogen 
 remaining uncombined. The temperature of the oxy- 
 hydrogen flame is much lower than this. The temperature 
 of an electric arc between carbon poles is estimated at about 
 
 1 All temperatures are given in Centigrade degrees. 
 
26 HYDROGEN 
 
 3600, while the temperature of an open hearth steel furnace 
 is only 1500 to 1700. 
 
 Platinum, which has a melting point of 1755, melts 
 readily in the oxyhydrogen flame. Iron takes fire and burns, 
 throwing off brilliant sparks. 
 
 The flame alone gives very little light, but if a piece of 
 lime is placed in the flame it glows intensely and gives a 
 very brilliant light, called variously the "lime light," 
 "calcium light'' or "Drummond light." This is often 
 used for illumination in stereopticons but has been mostly 
 replaced by electric lights. The oxyhydrogen blowpipe 
 is used to melt together the edges of the leaden plates used 
 in making the chambers for the manufacture of sulfuric 
 acid. A similar blowpipe in which acetylene is used in 
 place of hydrogen and which gives a still hotter flame is now 
 extensively used for cutting steel or iron bars or plates. 
 The flame will melt its way through the metal very rapidly. 
 The blowpipe is also used in blacksmith shops and garages 
 for welding. 
 
 SUMMARY 
 
 Water may be decomposed by red-hot iron, giving mag- 
 netic oxide of iron and hydrogen. 
 
 Sodium and water give hydrogen and sodium hydroxide. 
 Potassium gives hydrogen and potassium hydroxide. 
 
 Acids, especially hydrochloric acid and sulfuric acid, 
 contain hydrogen which can be displaced by zinc or iron, 
 giving the salts, zinc chloride or sulfate or ferrous chloride 
 or sulfate. 
 
 Hydrogen is the lightest gas known. It may be lique- 
 fied and frozen at very low temperatures. It burns in air, 
 forming water. 
 
 Hydrogen occurs in water, acids, organic compounds and, 
 in minute quantities, in air. 
 
 Hydrogen reduces hot copper oxide to metallic copper. 
 
SUMMARY. HYDROGEN 27 
 
 It also reduces magnetic oxide of iron but the reaction is 
 reversible. 
 
 The union of oxygen and hydrogen is also reversible at 
 high temperatures and this limits the temperature of the 
 oxyhydrogen blowpipe. 
 
 The oxyacetylene flame is used for cutting steel and iron. 
 
 EXERCISES 
 
 1. One part by weight of hydrogen requires eight parts by weight 
 of oxygen for its combustion. How many liters of oxygen will be 
 required to burn a liter of hydrogen? How many liters of air? 
 Suggestion: Find first from the text the weight of a liter of hydro- 
 gen and of a liter of oxygen and the proportion of oxygen in the 
 air. Answers are to be given in round numbers or with two or 
 three significant figures, not with long decimals. 
 
 2. Sugar contains 44 per cent of carbon and 6.7 per cent of 
 hydrogen. How many grams of oxygen will be required to burn a 
 pound of sugar? How many liters of air? 
 
 3. How many grams of water and how many grams of carbon 
 dioxide will be formed by burning a pound of sugar? 
 
 4. From a mixture of liquid air and liquid hydrogen in what 
 order would the elements distil away? What would be the effect 
 of applying a flame to such a mixture? 
 
 5. What is the "Hydione" method of preparing hydrogen? 
 
 6. How is the hydrogen to fill Zeppelins prepared? 
 
CHAPTER IV 
 
 WEIGHTS AND MEASURES, MOLECULAR THEORY, 
 LAWS OF GASES 
 
 Weights and Measures in Scientific Use. For scientific 
 purposes the weights and measures of the metric system are 
 used, almost exclusively. 
 
 The measure of length is the meter, which is 39.3709 
 inches. Its subdivisions are decimeters, centimeters, and 
 millimeters, which are, respectively, tenths, hundredths 
 and thousandths of a meter. 
 
 The measure of weight is the gram, which is almost, but 
 not exactly, the weight of one cubic centimeter of water 
 taken at its maximum density (i.e., at 4). Its most com- 
 mon subdivision is the milligram, one-thousandth of a 
 gram, and the most common multiple is the kilogram, which 
 is one thousand grams. 
 
 The measure of volume is the liter ; which is the volume 
 occupied by one kilogram of water at its maximum density. 
 It is almost, but not exactly, one cubic decimeter. The 
 most common division is the cubic centimeter, one-thou- 
 sandth of a liter. 
 
 The following approximate equivalents are sometimes 
 convenient : 
 
 One meter is a little more than a yard (39.3709 inches). 
 
 One millimeter is almost ^5 of an inch (0.03937 inch). 
 
 One gram is about 15 grains (15.432 grains). 
 
 One kilogram is about 2J pounds (2.2046 pounds). 
 
 One liter is a little less than a quart (0.8836 quart). 
 
 One cubic centimeter is about one-thirtieth of a fluid 
 ounce. 
 
 28 
 
MOLECULAR THEORY 29 
 
 Molecular Theory. Very many and very diverse phe- 
 nomena are most easily explained by supposing that all 
 material objects are composed of very minute particles, 
 which are called molecules. 1 In a proper scientific treat- 
 ment the various facts and relations which have led to this 
 conclusion should be presented as a foundation for this 
 theory. Practically, however, these facts are much more 
 easily understood and remembered if they are presented in 
 connection with the theory, and the theory is so useful that 
 it seems best to give an outline of it here. No student 
 should, however, be satisfied to accept the theory on the 
 authority of a book or a teacher. On the contrary, the 
 student should hold the theory in the earlier months of his 
 study as something which has not been fully demonstrated, 
 and he should return to the theory over and over again as 
 new facts related to it are learned. 
 
 Solids retain their shape unless they are subjected to 
 some force great enough to bend or break them. Liquids 
 flow to the bottom of the vessels which contain them and 
 present at their top a level surface, which is sharply sepa- 
 rated from the gas or vapor above. Gases give no similar 
 surface but completely fill the vessel in which they are con- 
 tained. It seems evident from these properties that the 
 molecules of solids are held rather rigidly in position by 
 attractive forces between them. The molecules of liquids, 
 on the other hand, must glide or slip easily over each other 
 but are still held together by attractive forces. The first 
 and most natural opinion about gases was that the molecules 
 repel each other and separate for this reason. Another 
 explanation of this property is given below. 
 
 Diffusion of Gases. If the mouths of two narrow cylin- 
 ders containing air and hydrogen are brought together 
 
 1 According to the molecular theory molecules are the smallest particles 
 of any substance which can exist alone. Atoms are the smallest particles 
 of an element. They may be identical with the molecules of the free 
 element but usually are not (p. 136). 
 
30 
 
 WEIGHTS AND MEASURES 
 
 (Fig. 12) with the cylinder containing hydrogen above, 
 
 although the air is 14)^ times as heavy as the hydrogen, 
 
 some of the air will make its way upward into the hydrogen 
 and some of the hydrogen will make its way 
 downward into the air. This can be shown by 
 testing the gas in each cylinder with a flame. 
 Each cylinder will be found to contain an ex- 
 plosive mixture of air and hydrogen. Such a 
 process by which two gases or two liquids in 
 contact with each other mix is called diffusion. 
 Two liquids in contact will diffuse into each 
 other only when they are mutually soluble. 
 Two gases in contact will always diffuse, even 
 though the density of one of the gases be one 
 hundred times that of the other and no matter 
 how different or how in- 
 soluble in each other the 
 same substances may be 
 
 when they are in the liquid state. 
 If a cylinder of porous porcelain, with 
 
 openings so fine that pressure will cause 
 
 a gas to pass through them only very 
 
 slowly, is fitted with a rubber stopper 
 
 and connected with a bulb and bent 
 
 tube containing water as shown in Fig. 
 
 13, when a beaker filled with hydrogen 
 
 is placed over the cylinder the pressure 
 
 within will suddenly increase and force 
 
 water out of the tube in a jet. This 
 
 shows that hydrogen passes through 
 
 the walls of the cylinder to the interior. 
 
 It can be shown in this case, also, that 
 
 some air passes out through the walls 
 
 of the cylinder but the light hydrogen passes through the 
 
 fine openings very much more rapidly than the air. 
 
 FIG. 12. 
 
KINETIC THEORY OF GASES 
 
 Kinetic Theory of Gases. Gases, under ordinary c 
 ditions, are much less dense than solids or liquids. If 
 quart of water is changed to steam at atmospheric pressure, 
 there will be more than 1600 quarts of steam. It seems 
 improbable that the particles (molecules) of water increase 
 very much in size when the water is changed to steam. 
 If they do not, it must be that the steam has 1600, or more, 
 quarts of empty space for every quart of space actually 
 filled by molecules of water. This and other facts, some 
 of which will be given below, have led to the proposal and 
 development of the kinetic theory of gases, a theory which 
 explains many of the properties of gases by considering the 
 motions of the molecules. According to this theory the 
 molecules of a gas are kept apart because they are in very 
 rapid motion and because they rebound like elastic balls 
 whenever they hit one another or when they hit any solid 
 substance. The pressure exerted by a gas is evidently 
 caused by the bombardment of any surface with which 
 it is in contact by the molecules of the gas, just as pressure 
 would be exerted on a wall if a large number of elastic balls 
 were constantly thrown against it. 
 
 It is evident that according to this theory a gas expands 
 and completely fills any space which is given to it, not be- 
 cause of any repulsion between the molecules of the gas, 
 but because the motion of the molecules causes them to 
 fly out and fill any space at the side, if there is no wall to 
 restrain them. Indeed, many facts which have been learned 
 about gases indicate that when the molecules come close 
 together in their collisions there is an attraction between 
 them and that it is largely for this reason that most gases 
 do not obey exactly the laws of temperature and pressure 
 which are given in the latter part of this chapter. 
 
 The kinetic theory gives a very satisfactory explanation 
 of the diffusion of gases. When two gases are brought into 
 contact the moving molecules of one can readily shoot into 
 
WEIGHTS AND MEASURES 
 
 spaces between the molecules of the other and some of 
 molecules which shoot in will rebound after collision in 
 a manner as to fly further into the mass of the, other 
 gas. As the velocities of the molecules are very great the 
 gases will rapidly become mixed at their surfaces and grad- 
 ually throughout their whole mass. This process of diffu- 
 sion can be made apparent to the eye by putting a little 
 bromine in the bottom of a tall cylinder containing air. 
 Bromine vapor is colored and its gradual diffusion into the 
 air can easily be seen. 
 
 Collision Between Elastic Bodies. When two elastic 
 balls of the same size, going in opposite directions, meet 
 
 FIG. 14. 
 
 squarely, each rebounds with a velocity such that its 
 energy is the same as that of the other ball before contact. 
 Thus if two balls, A and B, have the same weight, if A is 
 raised a certain distance on one side (Fig. 14) and allowed 
 to fall it will stop almost completely on hitting 5, while the 
 latter will rise to almost the same distance on the other side. 
 If A is raised through an arc of 60 while B is raised only 30, 
 
NUMBER OF MOLECULES 33 
 
 after impact B will recoil nearly 60 while A will recoil 
 only about 30. If one. ball is heavier than the other, the 
 lighter ball will recoil with a greater velocity than that 
 of the heavy one when struck by it and the heavy ball 
 will have a smaller velocity than that of the lighter one if 
 the heavy ball is at rest and is struck by the lighter one. 
 
 Number of Molecules in the Same Volume of Different 
 Gases. Applying these principles to the consideration 
 of a mixed gas containing molecules of different weights, 
 it has been shown that the average velocities of the mole- 
 cules of the different constituents of such a gas must vary 
 inversely as the square roots of the weights of the mole- 
 cules. Also, the greater velocity of the lighter molecules 
 will keep them just as far apart and give the same pressure 
 as the slower velocity of the heavy ones. It also follows 
 that when the pressure is the same there must be the same num- 
 ber of molecules in a given volume of a light gas as in the same 
 volume of a heavy one. The differences in the velocities of 
 light and heavy molecules explain very satisfactorily the 
 diffusion of hydrogen through the porous wall, and that 
 experiment, in turn, gives strong support for the kinetic 
 theory. Molecules of oxygen, according to the theory, 
 must be about 16 times as heavy as molecules of hydrogen 
 and will, therefore, fly only one-fourth as fast. As there are 
 the same number of molecules on both sides, four times as 
 many of the fast hydrogen molecules will hit the small 
 openings and fly through them in a given time as of the 
 slower oxygen molecules. 
 
 Number of Molecules in One Cubic Centimeter of a Gas. 
 Several different methods have been found for estimating 
 the number of molecules in a cubic centimeter of a gas 
 under standard conditions. The most accurate of these 
 indicate that the number is about 2.71 X 10 19 (or 27,100,- 
 000, 000, 000, 000, 000). The average velocity of hydro- 
 gen molecules at ordinary temperatures is somewhat more 
 
34 
 
 WEIGHTS AND MEASURES 
 
 than a mile a second. That of oxygen molecules is a little 
 more than a quarter of a mile a second. 
 
 The important conclusion that there are the same num- 
 ber of molecules in equal volumes of different gases under 
 the same conditions of temperature and pressure was 
 TEMPERA- VOLUME reached from a consideration 
 of chemical facts many years 
 373 cc. b e f ore tne kinetic theory of 
 gases was proposed. We shall 
 283 cc. have occasion to come back to 
 
 ABSOLUTE 
 TEMPERATURE 
 
 373 C 
 
 283 
 273 
 
 100 C 
 
 10 C 
 
 173 C 
 
 -100 -- 173cc. 
 
 73< 
 
 O c 
 
 -200< 
 
 -273 
 
 FIG. 15. 
 
 "I- 273 cc. this later (P- 135). 
 
 Temperature. Tempera- 
 tures are measured, for scien- 
 tific uses, with the Centigrade 
 thermometer. The freezing 
 point of water is taken as 
 and the boiling point of water, 
 under a pressure of one atmos- 
 phere (760 millimeters of mer- 
 cury), is taken as 100. 
 
 Absolute Zero. If the 
 pressure is kept constant, 
 hydrogen expands, when it is 
 heated, at the rate of 3/273 of 
 its volume at zero for each 
 degree. If it is cooled, it con- 
 tracts at the same rate, ^73 
 of its volume at zero, for each 
 degree. Evidently if it con- 
 tinued to contract at the same rate it would disappear at 
 a temperature of 273. It does not continue to contract 
 at the same rate,, of course, because it condenses to a liquid 
 at 252.5 and its expansion and contraction are then at 
 a very different rate. Other gases, however, expand and 
 contract, when heated or cooled, at almost the same rate 
 
 73 cc. 
 
LAW OF CHARLES 35 
 
 as hydrogen. This fact, that all gases expand and con- 
 tract at nearly the same rate for changes of temperature, has 
 made it seem appropriate to call 273 absolute zero, and 
 to call temperatures reckoned from that point absolute 
 temperatures. Before this temperature is reached all gases 
 become liquid, though the boiling point of helium is only 
 4.5 absolute. 
 
 The relations between ordinary and absolute tempera- 
 tures and between these and the volume of a gas which ex- 
 pands or contracts under constant pressure will be seen 
 from Fig. 15. 
 
 There are' many reasons for thinking that the absolute 
 zero is not merely a convenient fiction or theory based on 
 the conduct of gases. It seems to be an actual limit of 
 temperature which is a real starting point for many differ- 
 ent phenomena and which can never be reached by experi- 
 ment. The lowest temperature so far obtained is estimated 
 at about -270 or 3 absolute. 
 
 Law of Charles. The effect of a change in temperature 
 upon a gas under constant pressure is conveniently stated 
 in the law of Charles: The volume of a gas varies directly 
 as the absolute temperature. This may also be stated as a 
 proportion : 
 
 V : V : : T : T 
 
 where V and V ' represent two volumes of the same gas at the 
 temperatures T and T'. Thus 273 cc. (cubic centimeters) 
 of a gas at will become 283 cc. at 10 (= 273 + 10 or 
 283 absolute) or 263 cc. at -10 (= 263 absolute). 
 
 A very common problem in dealing with gases is to find 
 the volume which a gas measured at some given temperature 
 would occupy if cooled or heated to 0. Such a problem 
 may be easily solved by means of the proportion: 
 V Q :V::T<>(= 273):T( = t + 273) 
 
 V.-F 
 
36 
 
 WEIGHTS AND MEASURES 
 
 ,r 
 
 1 
 
 Such a formula should never be used mechanically, but 
 the student should merely remember that the volume 
 measured is to be multiplied by a fraction whose numerator 
 and denominator are the two absolute temperatures con- 
 cerned and that a gas expands when heated and contracts 
 when cooled. 
 
 Pressures. The pres- 
 sure of the air is most easily 
 and accurately determined 
 by measuring by means of 
 a mercury barometer the 
 height of the column of 
 mercury which will balance 
 it. For this reason the 
 pressure of a gas is usually 
 given in millimeters of mer- 
 cury, meaning the height, in 
 millimeters, of the column 
 of mercury which the 
 pressure of the gas will 
 sustain, or balance. On 
 1 the average, at sea-level, 
 
 the pressure of the at-- 
 ] mosphere will sustain a 
 
 Barometer ^J ^J ^J column of mercury 760 
 3 C mm. high, and this is taken 
 as the standard atmos- 
 pheric pressure. 
 Law of Boyle. When the pressure exerted on a gas is 
 increased or diminished the volume of the gas decreases or 
 increases in inverse proportion to the pressure. If we re- 
 call the kinetic theory it will be seen at once that this 
 should be true. If a gas is compressed to one-half its 
 volume, twice as many molecules must strike a square 
 centimeter of surface in a given time. This must cause 
 
 FIG. 16. 
 
VOLUME UNDER STANDARD CONDITIONS 37 
 
 twice the pressure on the surface, if the pressure is due 
 to the bombardment of the surface by the molecules. 
 
 The truth of this law may be demonstrated by means of 
 the apparatus shown in Fig. 16. If the mercury is at the 
 same level at both arms of the U-tube in A, it is evident 
 that the air in the graduated part will be at atmospheric 
 pressure (760 mm. at standard pressure). If mercury 
 is drawn out from the longer arm, as shown in B, till the 
 level in that arm is one-half the height of the barometer 
 column (380 mm.) lower than in the other arm, the pressure 
 in the graduated tube will be decreased to one-half of 
 an atmosphere and the volume of the air will be twice 
 as great as before. On the other hand, if mercury is poured 
 into the long arm till the level is as high as the length of 
 the barometer column (760 mm.) above the level in the 
 graduated tube, as shown in C, the pressure will be doubled 
 and the volume of the air will be decreased to one-half. 
 
 The law of Boyle may be conveniently stated by the 
 proportion : 
 
 V:V'::P':PorVP = V'P' or V = V'^ 
 
 For use in finding the volume of a gas under atmospheric 
 pressure when the volume under some other pressure is 
 known this becomes 
 
 Reduction of the Volume of a Gas to its Volume under 
 Standard Conditions. For the comparison of the properties 
 of different gases and for many other purposes it is con- 
 venient to have standard conditions of temperature and 
 pressure. The standard conditions which have been chosen 
 are Centigrade for temperature and 760 mm. for pressure. 
 When the volume of a gas is measured at some other 
 temperature and pressure, as is usually convenient, the 
 volume under standard conditions is found by multiply- 
 
38 
 
 WEIGHTS AND MEASURES 
 
 ing the observed volume by two fractions in which the 
 numerator and denominator of one are the two absolute 
 temperatures concerned and the numerator and denomi- 
 nator of the other are the two pressures. The formula is: 
 
 ' and 760 =: jari X ^T~ 
 
 To illustrate, suppose 15.2 cc. of gas are measured at 20 
 and under a pressure of 735 mm. The volume at and 
 
 735 273 
 760 mm. will be 15.2 X ^r^ X 
 
 As remarked of the formula for the law of Charles, such 
 a formula should never be used mechanically, but always 
 with a clear understanding of its meaning and remembering 
 that an increase in temperature innreaafis tfrfi 
 
 an increase in pressure decreases the volume of a gas. 
 ^Determination of the Weight of a Liter of a Gas. 
 Although gases are very light in comparison with solids 
 
 Toolrfxsrnfj 
 
 FlG. 17. 
 
 and liquids, they are heavy enough to be weighed with a 
 fair degree of accuracy on a sensitive balance. For this 
 purpose the bulb shown in Fig. 17 is exhausted till the 
 mercury in the manometer connected with it stands level 
 in the two arms. The stopcock is then closed and the 
 
WEIGHING A GAS 39 
 
 bulb is weighed. It is then filled with the gas to be ex- 
 amined, the temperature and barometer are read and the 
 bulb is weighed again. The increase in weight is, of course, 
 the weight of the gas contained in the bulb. The volume 
 of the bulb is determined by weighing it empty and full 
 of water. Having in this way determined the weight of a 
 known volume of the gas under a known temperature and 
 pressure, it is easy to calculate the volume which the same 
 gas would occupy at zero and atmospheric pressure, and' 
 from this the weight of one liter of the gas under standard 
 Conditions. 
 
 The process will be more clearly understood by calculating 
 the weight of a liter of air under standard conditions from 
 the following data: 
 
 Weight of bulb full of water 178 grams 
 
 Weight of empty bulb 64 grams 
 
 Weight of water : 114 grams 
 
 Hence the volume of the bulb is 114 cc. 
 
 Weight of the bulb full of air at 23 and 
 
 742 mm 64.0000 
 
 Weight of evacuated bulb 63 . 8668 
 
 Weight of air contained in the bulb at 
 23 and 742 mm 0. 1332 
 
 In solving this problem the volume of air in the bulb 
 is first reduced to its volume under standard conditions. 
 Remembering that a liter is 1000 cc. the calculation of the 
 weight of a liter of air is easy. 
 
 SUMMARY 
 
 The measures and weights most often used in scientific 
 work are the meter and millimeter, gram, milligram and 
 kilogram, and the liter and cubic centimeter. 
 
 According to the molecular theory substances are com- 
 posed of very small particles called molecules, which are 
 
40 WEIGHTS AND MEASURES 
 
 held together rather rigidly in solids, slide over each other 
 but are still held together by their attractions in liquids, 
 and are comparative^ far apart in gases. 
 ' Gases in contact always diffuse into each other. 
 
 The diffusion and pressure of gases are explained by the 
 kinetic theory, which supposes that the molecules of gases 
 are elastic bodies, in rapid motion. 
 
 Elastic bodies in collision exchange their energy of motion. 
 
 Equal volumes of different gases contain the same number 
 of molecules under the same conditions of temperature and 
 pressure. 
 
 One cubic centimeter of a gas under standard conditions 
 contains 2.71 X 10 19 molecules. 
 
 Temperatures are fixed with reference to the freezing 
 point and boiling point of water with 100 between. 
 
 An absolute temperature is the temperature measured 
 from 273 below zero. 
 
 The volume of a gas varies directly as the absolute tem- 
 perature (law of Charles) . 
 
 Pressures of gases are measured by the height in milli- 
 meters of the column of mercury the pressure will sustain. 
 Atmospheric pressure is taken as 760 mm. 
 
 The volume of a gas varies inversely as the pressure (law 
 of Boyle). 
 
 The two laws may be used in calculating volumes of 
 
 P' T 
 gases by means of the formula V = V X -p X TJT, 
 
 A gas may be weighed in a bulb which has been evacuated 
 and weighed empty and then full of the gas. 
 
 EXERCISES 
 
 1. What will be the volume under standard conditions of 
 37.5 cc. of a gas measured at 22 and a pressure of 735 mm? 
 
 2. What will be the volume at 25 and 730 mm. of 44.2 cc. of 
 gas measured at 27 and 770 mm.? 
 
EXERCISES. GASES 41 
 
 3. If one liter of air measured under standard conditions is 
 brought to a temperature of 12 and a pressure of 620 mm., what 
 will be its volume? 
 
 4. What will be the weight of a liter of air at 12 and 620 mm. 
 if it weighs 1.293 grams under standard conditions? What will 
 be the weight 01 a liter of hydrogen at the same temperature 
 and pressure, if it weighs 0.09 gram under standard conditions? 
 
 6. What will be the lifting power of a balloon having a volume 
 of 100 cubic meters and filled with hydrogen, if it is one mile 1 
 above sea-level and the temperature is 12 and the pressure 
 620 mm.? 
 
 6. Are there any other gases besides hydrogen which might be 
 used to fill balloons? 
 
 7. Since a gas expands and fills completely any space which is 
 given to it, why does not the atmosphere expand and fill the space 
 between it and the sun? 
 
 1 The height is not used in the calculation. It is given as the height 
 above sea-level where the pressure is approximately 620 mm. 
 
CHAPTER V 
 
 COMPOSITION OF WATER. LAWS OF COMPOSITION 
 BY WEIGHT. THE ATOMIC THEORY 
 
 Analysis. Synthesis. Two general methods, which 
 have been referred to briefly, are used in determining the 
 composition of substances: analysis and synthesis. In 
 analysis the substance is decomposed into its elements, or, 
 more often, the elements of which it is composed are con- 
 verted into other compounds which can be identified and 
 whose composition is known. The decomposition of oxide 
 of mercury into mercury and oxygen is the simplest kind of 
 analysis. The decomposition of steam by hot iron, with 
 the formation of magnetic oxide of iron and hydrogen 
 (p. 19) is also an analysis, when the composition of the oxide 
 of iron formed has been established as the same composition 
 as that of the oxide of iron formed by burning iron in 
 oxygen. 
 
 The composition of a substance is determined by synthe- 
 sis when the elements of which it is composed are put to- 
 gether to form it, as when hydrogen is burned in oxygen, 
 forming water. A synthesis of water may also be made by 
 passing hydrogen over heated copper oxide, giving copper 
 and water. An analysis or a synthesis may be qualitative, 
 giving simply the elements of which the substance is com- 
 posed, or quantitative, giving the proportion of each element 
 present. 
 
 Electrolysis of Sulfuric Acid. A very simple and a 
 roughly quantitative analysis of water can be made by 
 passing an electric current through dilute sulfuric acid, 
 
 42 
 
ELECTROLYSIS 
 
 using the apparatus shown in Fig. 18. Electrodes of plat- 
 inum in each arm of the U-tube are connected with the 
 poles of an electric battery by means of platinum wires 
 which pass through the glass. It can be shown that the 
 hydrogen of the sulfuric acid is carried toward the negative 
 electrode through the solution, while the rest of the acid, 
 composed of sulfur and oxygen, is carried 
 toward the positive electrode. At the 
 surface of the negative electrode the 
 hydrogen is liberated as a gas, while at 
 the positive electrode oxygen is liber- 
 ated. The volume of the hydrogen is 
 twice that of the oxygen. 
 
 Since a liter of hydrogen weighs 0.09 
 gram, and a liter of oxygen 1.429 grams, 
 the proportion by weight is 0.09 X 
 2 : 1.429 = 1 : 7.94, approximately. 
 The experiment is not suitable, how- 
 ever, for an accurate determination of 
 the composition of water. 
 
 As the hydrogen and oxygen are liber- 
 ated in the same proportion in which 
 they are found in water and as the total 
 amount of sulfuric acid in the solution 
 remains unchanged the experiment is 
 often spoken of as the decomposition of water by elec- 
 tricity. In consideration of the final result this is correct, 
 but it must not be overlooked that the motion of the 
 two parts of the sulfuric acid, hydrogen in one direction 
 and the sulfur and oxygen together in the other, arc an 
 essential part of the process of electrolysis. 
 
 The decomposition of a substance by passing a current of 
 electricity through it is called electrolysis. The poles con- 
 nected with the battery or dynamo which furnishes the 
 current of electricity are called electrodes. The substance 
 
 FIG. 18. 
 
44 
 
 COMPOSITION OF WATER 
 
 team 
 
 -18 
 
 =20 
 
 which is decomposed is called an electrolyte. The positive, 
 electrode, toward which the negative constituent of the 
 electrolyte moves, is called the anode. The negative elec- 
 trode, toward which the positive constituent of the elec- 
 trolyte moves is called the cathode. The hydrogen of the 
 sulfuric acid is positive and moves toward the 
 cathode, while the sulfur and oxygen together are 
 negative and move toward the anode. 
 
 Volumetric Composition of 
 Water. The proportion in 
 which hydrogen and oxygen 
 combine by volume to form 
 water may be more accurately 
 determined by mixing measured 
 quantities of oxygen and hydro- 
 gen in a graduated tube, ex- 
 ploding the mixture by means of 
 an electrical spark passed be- 
 tween platinum wires sealed 
 through the walls of the tube 
 and measuring the volume of 
 oxygen or hydrogen remaining 
 uncombined after the explosion. 
 Thus if we introduce 11 cc. of| 
 oxygen and 25 cc. of hydrogen 
 into the tube shown in Fig. 19' 
 and then explode the mixture, 
 it will be found that 3 cc. of hydrogen will remain uncom- 
 bined. Or if 15 cc. of oxygen are mixed with 18 cc. of 
 hydrogen, after explosion 6 cc. of oxygen will remain. The 
 gases must, of course, be measured at the same tempera- 
 ture and pressure in each case. If the tube is heated and 
 the water formed by the combination converted into 
 steam (Fig. 20), it is found that after taking account of the 
 increase in temperature and correcting the volume of the 
 
 FIG. 19. 
 
DUMAS'S EXPERIMENT 
 
 45 
 
 gas back to the original temperature there will be 25 cc. 
 of hydrogen and steam in the first case and 24 cc. of oxygen 
 and steam in the second case. A little study of these re- 
 sults will show that one volume of oxygen combines with 
 two volumes of hydrogen to give two volumes of steam. 
 This may be expressed graphically by the following diagram : 
 
 Oxygen 
 
 Hydrogen 
 
 Hydrogen 
 
 Steam 
 
 Steam 
 
 Composition of Water by Weight. The composition of 
 water by weight has been determined by reducing a weighed 
 amount of copper oxide to metallic copper by means of 
 hydrogen and collecting and weighing the water formed. 
 Thus hydrogen from the generator, F( Fig. 21), is passed 
 first through a series of tubes to purify and dry the gas and 
 then through the heated bulb, B, containing copper oxide, 
 where the hydrogen takes oxygen away from the oxide 
 and combines with it to form water. The water is collected 
 partly in a bulb, BI, just beyond the one containing copper 
 oxide, partly in tubes containing calcium chloride or other 
 substances which absorb vapor of the water, which might 
 otherwise escape. The loss in weight of the copper oxide 
 gives the weight of the oxygen, and the gain in weight of 
 the bulb and tubes in which the 1 water is collected gives 
 the weight of the water formed. With this apparatus the 
 hydrogen cannot be weighed directly, but its weight is 
 found by subtracting the weight of the oxygen from the 
 weight of the water. Dumas, the French chemist who 
 first used this method very carefully, did not succeed in 
 obtaining very accurate results with it. 
 
 By means of the apparatus shown in Fig. 22 it is possible 
 to weigh the hydrogen directly. The bulb A is filled with 
 copper oxide and, after exhausting it of air and closing the 
 
46 
 
 COMPOSITION OF WATER 
 
RATIO OF OXY.GEN TO HYDROGEN 
 
 47 
 
 stopcock, it is weighed. It is then placed in an air-bath so 
 that it can be heated, while the side tube B is cooled and 
 the tube C is connected with an apparatus giving pure, dry 
 hydrogen. As this passes over the hot copper oxide it is 
 oxidized to water and the latter is condensed in the tube B. 
 After a considerable amount of water has collected, the 
 apparatus is cooled and weighed again. Evidently the 
 gain in weight is the* weight of the hydrogen which has en- 
 tered the apparatus, as all of the oxygen of the copper oxide 
 
 B 
 
 FIG. 22. 
 
 remains within the apparatus either as unreduced copper 
 oxide or as water. By heating the apparatus and side tube 
 after connecting it with another bulb from which the air 
 has been exhausted, the water can be driven out and 
 collected. The difference between the weight of the bulb at 
 first and the weight after removing the water gives the 
 weight of the oxygen which has combined with the hydrogen. 
 The apparatus in which the water is collected may also be 
 weighed and the weight of the water formed determined. 
 The results of experiments by this method showed that the 
 ratio between the weights of hydrogen and oxygen in water 
 is H : O = 1 : 7.938. Determinations by other still more 
 accurate methods have given almost exactly the same 
 result. 
 
48 COMPOSITION OF WATER 
 
 Law of Constant Proportion. The more carefully the 
 experiments for determining the composition of water are 
 carried out the more exactly do the results agree with the 
 value for the ratio which has been stated above. This 
 illustrates the law of constant proportion, which has already 
 been given (p. 3). 
 
 Law of Combining Proportions. We may select for each 
 element some number which may always be used to repre- 
 sent the proportion of the element which enters into combi- 
 nation with other elements. This may be illustrated by 
 the following list of compounds: 1 
 
 Water Cuprous Oxide Cupric Sulfide Hydrogen Sulfide 
 
 H : : Cu Cu : S S : H 
 
 1:8 8 : 63.6 63.6 : 32 32:2 
 
 Hydrochloric Cupric Chloride Cupric Oxide Chlorine 
 Acid Dioxide 
 
 H:C1 Cl:Cu Cu : O : Cl 
 
 2:71 71: 63.6 63.6 : 16 16 : 17.75 
 
 These compounds are selected so that each contains one 
 element of the preceding and one of the following compound. 
 The table might be extended to contain a thousand com- 
 pounds and it would be found that in all compounds of 
 oxygen the amount combining with any other element 
 would be either 8 parts or 8 parts multiplied or divided by 
 some 'whole number. In the same way the amount of 
 hydrogen would be 1, 2, 3, 4, or more parts. The propor- 
 tion of copper entering into combination with other elements 
 is 63.6 parts for the cases given, and if the table were ex- 
 tended to contain other compounds of copper the amount 
 of that element combining with other elements would 
 always be 63.6 parts or some multiple or submultiple 
 of 63.6 parts. We might select the combining proportions, 
 
 1 For the sake of simplicity round numbers are used here and elsewhere. 
 
LAW OF MULTIPLE PROPORTIONS 49 
 
 H = 1, = 8, Cu = 63.6, S = 32, Cl = 71, and we could then 
 express the composition of every compound of these ele- 
 ments by these numbers or by multiples or submultiples 
 of these numbers. The values O = 16 and Cl = 35.5 are 
 practically used in place of 8 and 71. 
 
 The combining proportions which have been selected are 
 called atomic weights, because it is believed that they repre- 
 sent the relative weights of those smallest particles of the 
 elements, which are called atoms (p. 29). A table of 
 atomic weights is given on p. 164 and on the inside of the 
 first cover. 
 
 Hydrogen Peroxide. There is a second compound of 
 oxygen and hydrogen which contains a larger proportion 
 of oxygen than water does. The chemical name of water is 
 hydrogen oxide, but that name is very rarely used. Water 
 contains, in round numbers, 1 part of hydrogen to 8 parts 
 of oxygen. Hydrogen peroxide contains 1 part of hydrogen 
 to 16 parts of oxygen and the name given to it indicates 
 this composition, the prefix per meaning "more of." In 
 this case it means that the compound contains more oxygen 
 than common water. 
 
 Hydrogen peroxide, when pure, has a much higher 
 specific gravity than water; It is a colorless liquid and 
 it is very unstable. When pure it is liable to decompose 
 explosively into water and oxygen gas. 
 
 Nearly all explosions depend on the formation of a large 
 volume of gas from some solid or liquid substance. 
 
 A solution of hydrogen peroxide in water is more stable 
 than the pure compound and is used by dentists and others 
 as a germicide. Hydrogen peroxide is also a good bleaching 
 agent, especially for hair or silk. 
 
 Law of Multiple Proportions. The weight of oxygen 
 combined with 1 part by weight of hydrogen is exactly 
 twice as great in hydrogen peroxide as it is in water. The 
 relation may also be given by saying that there is twice as 
 
50 COMPOSITION OF WATER 
 
 much hydrogen in water, for 1 part of oxygen, as there 
 is in hydrogen peroxide. Stated in the form of a proportion 
 these are: 
 
 Water Hydrogen Peroxide 
 
 H: H : 
 
 1:8 1 : 16 
 
 or 2 : 16 2 : 32 
 
 or K : 1 He : 1 
 
 In many other cases two elements form two or more 
 compounds with each other and in every case relations 
 similar to this are found. This has given us the law of 
 multiple proportions: // two elements combine in different 
 proportions to form two different compounds, when we con- 
 sider a fixed amount of one element, the amounts of the other 
 element which combine with this amount will bear a simple 
 ratio to each other. 
 
 The following still more striking illustration of this law 
 may be given: 
 
 N 
 
 Nitrous oxide 14 
 
 Nitric oxide 14 
 
 Nitrogen trioxide 14 
 
 Nitrogen tetroxide 14 
 
 Nitrogen pentoxide 14 
 
 O N: O 
 
 8 or N 2 28 : 16 
 
 16 or NO 28 : 32 
 
 24 or N 2 3 28 : 48 
 
 32 or N 2 4 28 : 64 
 
 40 or N 2 5 28 : 80 
 
 A study of this law shows that it follows, necessarily, 
 from the law of combining proportions and that the law 
 of combining proportions is more general in its application 
 and more important. The law of multiple proportions is 
 simpler, however, and was discovered first. 
 
 The Atomic Theory. A number of reasons, have been 
 given for believing that gases and other substances consist 
 of very small particles, which are called molecules. The 
 law of multiple proportions was discovered by Dalton 
 about 1804 and in trying to find some reason for it he came 
 to the conclusion that the molecules of compounds must be 
 
SYMBOLS. FORMULAS 51 
 
 formed by the union of still smaller particles which he 
 called atoms of the elements. He supposed further that 
 each atom of an element weighs the same as every other 
 atom of the same element, but that the atoms of different 
 elements have different weights. He also ' supposed that 
 chemical combination always takes place between atoms. 
 If this theory is true, the laws of constant proportion, of 
 combining weights, and of multiple proportions follow 
 directly from it. Further than that, while we cannot 
 weigh an atom or molecule by any ordinary method, we 
 can determine the relative weights of the atoms by deter- 
 mining the composition by weight of the compounds which 
 they form. Thus if we can discover in some way (see p. 131) 
 that a molecule of water consists of two atoms of hydrogen 
 combined with one atom of oxygen, when we have deter- 
 mined that 1 part by weight of hydrogen combines with 
 8 parts by weight of oxygen to form water we can at once 
 say that an atom of oxygen is 16 tunes as heavy as an atom 
 of hydrogen. 1 Evidently the atomic weights, if we can 
 determine them by such a process, will be the most satis- 
 factory numbers to use as the combining proportions of the 
 elements, to express the composition of compounds. 
 
 Symbols, Formulas. It has been found convenient to use 
 the symbol of an element to stand for an atom of the 
 element. The number of atoms of each element in a mole- 
 cule of a compound is designated by a small figure placed 
 below the line and to the right of each symbol. Thus the 
 composition of a molecule of water is represented by the 
 formula H 2 O, which means that one molecule of water 
 contains two atoms of hydrogen and one atom of oxygen. 
 The formula might be a little clearer, perhaps, if it were 
 written H 2 0i, but the number one is always understood 
 
 1 Round numbers are used. The true value is 15.88. As a matter of 
 convenience, however, 16 has been selected as the atomic weight of oxygen 
 and the accurate atomic weight of hydrogen is 1.0077. 
 
52 COMPOSITION OF WATER 
 
 when no figure is used with a symbol. The formula of 
 hydrogen peroxide is H 2 O 2 , that of sulfur dioxide SO 2 , 
 that of carbon dioxide CO 2 , of phosphorus pentoxide 
 P2O 5 , of magnetic oxide of iron Fe 3 O4, of sodium hydroxide 
 NaOH. 
 
 The atomic weights, in round numbers, for the elements 
 most used in these formulas are: 
 
 H = 1, O = 16, S = 32, C = 12, P = 31, 
 
 Fe = 56, Na = 23. 
 
 Distinction between Parts by Weight and Atoms. If 
 we remember the atomic weight, the formula of a compound 
 tells us its composition by weight. Water, H 2 O, contains 
 two atoms of hydrogen for one atom of oxygen and it 
 contains 2 parts by weight of hydrogen for 16 parts 
 by weight of oxygen. It should be noticed, especially, 
 that the formula does not mean that water contains 2 parts of 
 hydrogen for 1 part of oxygen, though such a misinterpreta- 
 tion of the formula is often made' by beginners. Phosphorus 
 pentoxide, P 2 O5, contains twice 31, or 62 parts of phosphorus 
 combined with five times 16, or 80 parts of oxygen. Sodium 
 hydroxide, NaOH, contains 23 parts of sodium, 16 parts 
 of oxygen and 1 part of hydrogen. 
 
 Equations. Symbols and formulas may be combined 
 in equations which furnish a very concise statement of what 
 happens in chemical reactions. Thus the equation: 
 
 Na + H 2 O = NaOH + H 
 
 means that one atom of sodium acts on one molecule of 
 water to form one molecule of sodium hydroxide, NaOH, 
 and one atom of hydrogen. It means, also, that 23 parts 
 of sodium with 18 parts of water give 40 parts of sodium 
 hydroxide and 1 part of hydrogen. 
 The equation: 
 
 3Fe + 4O = Fe 3 O* 
 
EQUATIONS. CALCULATIONS 53 
 
 means that three atoms of iron combine with four atoms 
 of oxygen to form one molecule of the magnetic oxide; also 
 that 3 X 56 = 168 parts of iron combine with 4 X 16 = 64 
 parts of oxygen to form 168 + 64 = 232 parts of the 
 magnetic oxide of iron. 
 
 Writing Equations. Students often make the mistake 
 of memorizing equations. This ought never to be done. 
 The first step should always be to learn the formulas of the 
 compounds which act on each other and the formulas of 
 the products formed in the reaction. With these formulas 
 as a starting point, the equation should be deduced logically. 
 
 Thus when steam, H 2 O, is passed over heated iron, Fe, 
 magnetic oxide of iron, Fe 3 O4, and hydrogen, H, are formed. 
 With this starting point we write: 
 
 Fe + H 2 O-Fe 3 4 + H., 1 
 
 On looking at these formulas we see that there are four 
 atoms of oxygen on the right side and only one on the 
 left ; there are also three atoms of iron on the right and only 
 one on the left. We must have, therefore: 
 
 3Fe + 4H 2 0-Fe 3 4 + H 2 
 
 Looking again at this we see that there are 4 X 2 or 8 
 atoms of hydrogen on the left and we must have the same 
 number on the right. The equation then becomes: 
 
 3Fe + 4H 2 O = Fe 3 O 4 + 4H 2 
 
 If we wish to express the fact that the reaction is re- 
 versible, we may use two arrows in place of the sign of 
 equality. Thus: 
 
 3Fe + 4H 2 0<FFe 3 O 4 + 4H 2 
 
 Calculations. As a formula indicates the composition 
 of a compound in terms of the atomic weights the percentage 
 composition can always be calculated from it. The formula 
 
 1 The reason for writing the formula of hydrogen H? and that of oxygen 
 O 2 will be given later (p. 136). 
 
54 COMPOSITION OF WATER 
 
 of potassium chlorate is KC10 3 . The composition by 
 
 weight is, from this formula: 
 
 Potassium, K = 39 parts 
 Chlorine, Cl = 35.5 parts 
 Oxygen, 3 = 3 X 16 or 48 parts 
 Total = 122.5 parts 
 
 The percentage of oxygen must be 100 X 100 = 39.2 
 
 per cent. 
 
 If we have the equation for any reaction and know the 
 amount of any substance used or formed in the reaction 
 we can calculate the amounts of every other substance 
 involved. 
 
 Suppose we have 25 grams of iron, Fe, and wish to know 
 how much water it will decompose and what weights of 
 magnetic oxide and of hydrogen will be formed. 
 
 The first step in solving any problem of this sort should 
 be to write the equation for the reaction : 
 
 3Fe + 4H 2 O = Fe 3 O 4 + 4H 2 
 
 Next write the formulas of the two substances to be 
 considered, in the form of two identical ratios, using the 
 number which is used before each formula in the equation 
 before the same formula in the ratios. Beneath the for- 
 mulas of the first ratio write the corresponding atomic or 
 molecular weights. Beneath the formula of the known 
 substance in the second ratio write the amount of that sub- 
 stance stated in the problem. Beneath the other formula 
 in the second ratio write "x". Finally solve the proportion 
 for x. 
 
 Thus to calculate the amount of water, above: 
 
 3Fe:4H 2 O ::3Fe:4H 2 
 3 X 56:4 X 18:: 25 : x 
 168:72::25 : x 
 
 x = = 10.71 grams of water 
 
 .LOo 
 
SUMMARY. WATER, ATOMIC THEORY 55 
 
 In a similar manner the amounts of magnetic oxide and 
 of hydrogen may be calculated. 
 
 The student should also notice that since 168 parts of 
 iron will decompose 72 parts of water, the amount of water 
 
 is found by multiplying the weight of iron, 25 grams, by 
 72 
 
 Problems of this sort may be solved by using a frac- 
 
 tion in this way without writing the formal proportion. 
 The fraction is less likely to be used mechanically than is the 
 proportion. 
 
 SUMMARY 
 
 The composition of substances is determined by analysis, 
 the separation into the elements of which they are com- 
 posed; or by synthesis, preparation by putting the elements 
 together. 
 
 A simple analysis of water may be made by the electroly- 
 sis of dilute sulfuric acid, giving two volumes of hydrogen 
 to one of oxygen. 
 
 In electrolysis the poles of the battery put into the elec- 
 trolyte are called electrodes. 
 
 The positive electrode is the anode, the negative is the 
 cathode. 
 
 A mixture of measured volumes of oxygen and hydrogen 
 may be exploded, and by measuring the volume of the gas 
 remaining an accurate synthesis of water can be made. 
 
 By passing hydrogen into an evacuated apparatus con- 
 taining copper oxide, and afterward removing the water 
 formed, a synthesis of water by weight can be effected. 
 The gain in weight gives the weight of the hydrogen; the 
 loss after removing the water gives the weight of the oxygen. 
 The water may also be collected and weighed. 
 
 A pure compound always has exactly the same composi- 
 tion by weight (law of constant proportion). 
 
 A number may be selected for each element which repre- 
 
56 COMPOSITION OF WATER 
 
 sents the proportion by weight in which the element enters 
 into combination with any other element (law of combining 
 proportions). The most convenient numbers for this pur- 
 pose are the atomic weights of the elements. 
 
 Hydrogen peroxide is a second compound of hydrogen and 
 oxygen. It contains twice as much oxygen for a given 
 weight of hydrogen as water does. 
 
 Hydrogen peroxide and water illustrate the law that 
 whenever two elements form two compounds having a 
 different composition the amounts of one which combine 
 with a fixed amount of the other will bear a simple ratio 
 to each other (law of multiple proportions). 
 
 The atomic theory gives a simple and satisfactory expla- 
 nation of the laws of constant proportions, combining pro- 
 portions and multiple proportions. 
 
 Symbols designate: 
 
 (a) An element. H stands for hydrogen; Fe for iron. 
 
 (6) An atom of an element. stands for an atom of 
 oxygen. 
 
 (c) A gram atom of the element. Fe stands for 56 grams 
 of iron. 
 
 Formulas designate: 
 
 (a) A compound. H 2 O stands for water; KC1O 3 for 
 potassium chlorate. 
 
 (6) A molecule of a compound. H 2 O 2 stands for a mole- 
 cule of hydrogen peroxide; O 2 stands for a molecule of 
 oxygen. 
 
 (c) A gram-molecule of a compound. H 2 O stands for 
 18 grams of water. 
 
 (d) The composition by weight of a compound. H 2 O 
 indicates that water is composed of 2 parts of hydrogen 
 combined with 16 parts of oxygen. 
 
 Equations are used to represent: 
 
 (a) The substances which react and the substances 
 formed in a chemical reaction. 
 
EXERCISES. WATER, ATOMIC THEORY 57 
 
 (6) The atomic composition of each substance involved. 
 (c) The quantities of the substances. 
 Equations should be true : 
 (a) Chemically. 
 (6) Algebraically. 
 
 Formulas and equations may be used as the basis for 
 a great variety of calculations. 
 
 EXERCISES 
 
 1. Write equations to represent the burning of sulfur, charcoal, 
 phosphorus and iron. The formulas of the compounds formed 
 are S0 2 , C0 2 , P 2 5 and Fe 3 4 . 
 
 2. Write the equations for the decomposition of oxide of mer- 
 cury, HgO, and of potassium chlorate, KClOs. 
 
 3. At a high temperature manganese dioxide, Mn0 2 , may be 
 decomposed into oxygen, 2 , and another oxide, Mn 3 4 . Write 
 the equation for the reaction. 
 
 4. Write the equations for the action of hydrochloric acid, HC1, 
 and of sulfuric acid, H 2 S0 4 , on zinc and on iron. The compounds 
 formed are zinc chloride, ZnCl 2 , ferrous chloride, FeCU, zinc 
 sulfate, ZnS0 4 , and ferrous sulfate, FeS0 4 . 
 
 5. Using the atomic weights, H = 1, 0. = 16, Cu = 63.6, 
 S = 32, Cl = 35.5, write the formulas used in illustrating the law 
 of combining proportions. 
 
 6. Calculate the percentage composition of mercuric oxide, HgO 
 (Hg = 200), potassium chlorate, KC10 3 (K = 39, Cl = 35.5), 
 and sulfuric acid, H 2 SO 4 (S = 32). 
 
 7. If a gasometer holds 20 liters, how much potassium chlorate 
 will be required to give enough oxygen to fill it? One liter of 
 oxygen weighs 1.429 grams. 
 
 8. How much zinc and how much sulfuric acid will be required 
 to fill the same gasometer with hydrogen? One liter of hydrogen 
 weighs 0.09 gram. 
 
 9. Air contains 21 per cent by volume of oxygen. How many 
 liters of air will be required to burn one pound (453 grams) of 
 sulfur (S = 32)? How many liters will be required to burn one 
 pound of charcoal, supposing it to be pure carbon (C = 12)? 
 
CHAPTER VI 
 PROPERTIES AND USES OF WATER. VAPOR PRESSURE 
 
 Freezing Water. When ice is heated it begins to melt 
 at and if the ice and the water formed (solid and liquid 
 phases) are kept in intimate contact the temperature of 
 the mixture will remain at till all the ice is melted. If 
 water containing no ice is cooled, its temperature may fall 
 considerably below the freezing point before any ice begins 
 to form. Water in this condition is said to be supercooled. 
 It is in an unstable condition and if some ice is thrown into 
 this supercooled water part of the water in contact with 
 the ice will freeze and in freezing will give out heat to the 
 rest of the water until the temperature of the whole rises 
 to zero. Ice and water cannot be in contact with each 
 other at any other temperature. 
 
 If two dishes, one containing a kilogram of ice and the 
 other containing a kilogram of water at the freezing point, 
 are placed over two gas burners in such a manner that each 
 receives heat at the same rate, when the ice is just melted 
 and the water in contact with it is still at zero the water 
 in the other dish will be found to have reached a tempera- 
 ture of 80. In other words, it requires as much heat to 
 melt a pound of ice without changing its temperature as it 
 does to raise the temperature of a pound of water 80 or 
 to raise the temperature of 80 pounds of water one degree. 
 The latent heat of fusion of ice is said to be 80 calories 
 (accurately 79.63 calories). 
 
 A calorie is the heat required to raise the temperature of 
 one kilogram of water one degree. The small calorie, the 
 
 58 
 
CHANGES OF VOLUME OF WATER 59 
 
 heat required to raise the temperature of a gram of water 
 one degree, is also often used. 
 
 Boiling of Water. If water is heated to 100 under at- 
 mospheric pressure (760 mm.), bubbles of steam will form 
 and rise through the water, and if some means is provided to 
 secure the continual formation of minute bubbles in the 
 water, the water will continue to boil and the temperature 
 will remain at 100 till all of the water has been converted 
 into steam. If no provision for starting bubbles of steam 
 is made, the water may become superheated to a temperature 
 considerably above 100, but if bubbles of gas are introduced 
 in superheated water the water will boil rapidly and the 
 temperature will fall to 100, the true boiling point. 
 
 If' two flasks, one containing a kilogram of water at 
 and the- other containing a kilogram of water at 100, are 
 placed over two burners so that the same quantity of 
 heat enters each, the water at 100 will boil away without 
 changing its temperature, but when the temperature of the 
 water in the first flask has risen to 100 d less than one-fifth 
 of the water in the second flask will have been changed to 
 steam. If another flask at is placed over the first burner 
 and left till its temperature reaches 100 and this is repeated 
 until all the water in the second flask is converted into 
 steam, it will be found that the water in five flasks has been 
 heated from to 100 and that the water in the sixth 
 flask will be at a temperature of 36. The latent heat of 
 vaporization of water at 100 is said to be 536 calories (accu- 
 rately 536.6).' 
 
 Changes of Volume of Water for Changes of Tempera- 
 ture. With very few exceptions, substances of all kinds, 
 whether solids, liquids or gases, expand when heated and 
 contract when cooled. A number of liquids expand on 
 solidifying, although few, if any, expand as much as water 
 when it freezes. But water is almost, or quite, alone in 
 expanding on cooling while the temperature is still several 
 
60 
 
 PROPERTIES AND USES OF WATER 
 
 degrees above the freezing point. When water is cooled it 
 contracts as other substances do till a temperature of 
 4 above the freezing point is reached. If cooled further, 
 it begins to expand and 8000 cc. of water at 4 will fill a 
 volume of 8001 cc. at 0. On freezing, the same amount of 
 water will give 8727 cc. of ice. The expansion of water 
 on freezing is familiar in the bursting of a pitcher or a 
 bottle filled with water, ink or other aqueous liquid, when 
 the liquid freezes. While the increase of volume from 4 
 to is small, it is very important practically, since this 
 property of water protects the ocean and other large 
 bodies of water from cooling to the freezing 
 point, because the lighter, cooler water 
 floats on the surface and is a poor con- 
 ductor of heat. As long as the tempera- 
 ture of a body of water is above 4 the 
 colder water on top falls and is replaced 
 by .warmer water from below, but at 4 
 this method of cooling (called cooling by 
 convection) ceases. The ice which finally 
 forms on the surface still further protects 
 the mass of water below from cooling below 
 a temperature of 4, because it is a very 
 poor conductor of heat and cannot be 
 cooled by convection. 
 
 Vapor Pressure of Water. If a drop of 
 water is introduced into the space above 
 the mercury in a barometer (Fig. 23), a little 
 of the water will evaporate and the vapor 
 of the water will exert a pressure on the mercury and cause 
 the mercury to fall about 20 mm. (^5 inch) if the tem- 
 perature is 22 (72 F.) . If the temperature is higher the 
 mercury will fall further; if it is lower it will fall less. For 
 any given temperature the pressure of the vapor is always 
 the same, provided only water and vapor are present in 
 
 FIG. 23. 
 
YAPOR PRESSURE 
 
 61 
 
 the space above the mercury. If a few drops of water are 
 introduced into air which has been carefully dried and 
 which is contained in a closed bottle (Fig. 24), some of the 
 water will evaporate and the vapor of the water will diffuse 
 into the air of the bottle and will increase the pressure 
 within the bottle, as shown by the manometer (A) con- 
 nected with it. This increase in pressure will be found to be 
 the same as the pressure of the 
 vapor in the barometer tube 
 above the mercury. This is be- 
 cause the molecules of water in 
 the form of vapor, which is 
 merely a liquid or solid con- 
 verted into a gas, shoot in 
 between the molecules of air 
 (p. 31) and ultimately there are 
 just as many molecules of vapor 
 as there would be if the air were 
 not present. 
 
 In any mixture of gases each 
 gas exerts the same pressure as 
 it would exert if it were present 
 alone in the same space (Henry's 
 law). 
 
 If the tube containing the 
 mercury and water is heated to 100 while the pressure of the 
 air is 760 mm., the mercury within and without the tube 
 will be at the same level. In other words, the boiling point 
 of water is that temperature at which the pressure of the 
 vapor of water is equal to the pressure of the atmosphere. 
 
 Natural Waters. Rain water which falls in the open 
 country, after it has been raining some time so that the 
 substances which would contaminate the water have been 
 washed out of the air, is very nearly pure. As soon as the 
 water reaches the earth, however, it begins to take up 
 
 FIG. 24. 
 
62 PROPERTIES AND USES OF WATER 
 
 various substances. Carbonic acid, H 2 CO 3 , formed by 
 the oxidation of vegetable and animal substances (" organic 
 matter") in the soil, is taken up by the water, and if the soil 
 also contains small particles of limestone, or calcium 
 carbonate, CaCO 3 , as is usually the c&se, this dissolves in 
 the carbonic acid of the water and makes it "hard." 
 Such a water requires more soap than soft water to produce 
 a lather because the soap is at first used up in producing 
 a curdy precipitate with the calcium of the calcium car- 
 bonate. On boiling the water the carbonic acid, H 2 CO 3 , 
 which holds the calcium carbonate in solution, breaks up 
 into carbon dioxide, CO 2 , and water and the calcium car- 
 bonate is precipitated. This causes a scale in teakettles 
 and boilers. 
 
 Equilibrium between a Liquid and its Vapor. Kinetic 
 Theory. In discussing the kinetic theory of gases (p. 31) 
 it was pointed out that while the molecules of gases are so 
 far apart and in such rapid motion that they exert little 
 attraction upon each other, the molecules of a liquid are 
 held together by attractive forces. It does not follow from 
 this that the molecules of a liquid are not also in motion. 
 Indeed we cannot well explain the escape of a vapor from the 
 surface of a liquid without supposing that the molecules 
 of a liquid are in motion and that occasionally the motion 
 of an individual molecule is sufficiently rapid so that it 
 tears itself away from its fellows and shoots out into the 
 space above. Such a molecule then becomes a gaseous 
 molecule and it will continue straight on until it meets with 
 some obstacle. On the other hand, an occasional molecule 
 of the vapor will be moving toward the liquid surface and 
 when it reaches this it may be held by the attraction of 
 the molecules of the liquid and become a part of the liquid 
 once more. At any given temperature the number of 
 molecules leaving the liquid and the number returning to 
 it must be the same, when a state of equilibrium is reached. 
 
TABLE OF VAPOR PRESSURE 
 
 63 
 
 If the volume of the vapor above the liquid is increased 
 by enlarging the space filled by it, as, for instance, by rais- 
 ing the piston in Fig. 25, the pressure of the vapor will be 
 decreased in accordance with the laws of gases. Immedi- 
 
 VAPOR PRESSURE OF ICE AND WATER 
 
 TEMPERA- 
 TURES 
 
 PRESSURE IN 
 MILLIMETERS 
 OP MERCURY 
 
 TEMPERA- 
 TURES 
 
 PRESSURE 
 
 IN MM. 
 
 TEMPERA- 
 TURES 
 
 PRESSURE 
 
 IN MM. 
 
 -10 
 
 2.0 
 
 27 
 
 26.5 
 
 100.5 
 
 773.7 
 
 - 5 
 
 3.0 
 
 28 
 
 28.1 
 
 101.0 
 
 787.6 
 
 - 2 
 
 3.9 
 
 29 
 
 29.8 
 
 
 
 - 1 
 
 4.2 
 
 30 
 
 31.6 
 
 
 
 
 
 4.6 
 
 31 
 
 33.4 
 
 
 
 1 
 
 4.9 
 
 32 
 
 35.4 
 
 
 
 2 
 3 
 
 5.3 
 
 5.7 
 
 33 
 34 
 
 37.4 
 39.6 
 
 
 PRESSURE IN 
 
 4 
 
 6.1 
 
 35 
 
 41.9 
 
 
 ATMOSPHERES 
 
 5 
 
 6.5 
 
 40 
 
 55.0 
 
 
 
 
 6 
 
 7.0 
 
 50 
 
 92.2 
 
 111.7 
 
 1.5 
 
 7 
 
 7.5 
 
 60 
 
 149.2 
 
 120.6 
 
 2 
 
 8 
 
 8.0 
 
 70 
 
 233.8 
 
 127.8 
 
 2.5 
 
 9 
 
 8.6 
 
 80 
 
 355.5 
 
 133.9 
 
 3 
 
 10 
 
 9.2 
 
 90 
 
 526.0 
 
 144.0 
 
 4 
 
 11 
 
 9.8 
 
 95 
 
 634.0 
 
 159.2 
 
 6 
 
 12 
 
 10.5 
 
 96 
 
 657.7 
 
 170.8 
 
 8 
 
 13 
 
 11.2 
 
 97 
 
 682.1 
 
 180 . 3 
 
 10 
 
 14 
 
 11.9 
 
 98 
 
 707.3 
 
 188 . 4 
 
 12 
 
 15 
 
 12.7 
 
 99 
 
 733.2 
 
 195 . 5 
 
 14 
 
 16 
 
 13.6 
 
 99.1 
 
 735.9 
 
 201.9 
 
 16 
 
 17 
 
 14.5 
 
 99.2 
 
 738.5 
 
 207.7 
 
 18 
 
 18 
 
 15.4 
 
 99.3 
 
 741.2 
 
 213.0 
 
 20 
 
 19 
 
 16.4 
 
 99.4 
 
 743.9 
 
 224.7 
 
 25 
 
 20 
 
 17.4 
 
 99.5 
 
 746.5 
 
 
 
 21 
 
 18.5 
 
 99.6 
 
 749.2 
 
 
 
 22 
 
 19.7 
 
 99.7 
 
 751. 9* 
 
 
 
 23 
 
 20.9 
 
 99.8 
 
 754.6 
 
 
 
 24 
 
 22.2 
 
 99.9 
 
 757.3 
 
 
 
 25 
 
 23.5 
 
 100.0 
 
 760.0 
 
 
 
 26 
 
 25.0 
 
 100.1 
 
 762.7 
 
 
 
 
 
 100.2 
 
 765.5 
 
 
 
 
 
 100.3 
 
 768.2 
 
 
 
 
 
 100 . 4 
 
 770.9 
 
 
 
64 PROPERTIES AND USES OF WATER 
 
 ately, however, a larger number of molecules will leave 
 the surface of the liquid than will return to it because 
 fewer strike the surface, and this will continue till the 
 pressure is the same as before. If the piston is forced 
 down and the pressure of the vapor is temporarily increased, 
 the opposite process will result. For each temperature 
 every liquid has a definite pressure for its vapor in 
 contact with the liquid. The same law applies to 
 solids which are volatile. Ice has a vapor pressure 
 which decreases as it grows colder. 
 
 The table on page 63 gives the vapor pressures 
 of ice and water at different temperatures. 
 
 Germs of Diseases Carried by Water. It often 
 happens that by means of sewage disease germs 
 coming from cases of typhoid fever and from other 
 diseases find their way into rivers or lakes or wells. 
 Persons who drink water that has been contam- 
 inated in this manner are liable to contract the 
 disease which is caused by these germs. The 
 FIG 25 ^ ac ^ er ^ a which are the exclusive cause of such 
 diseases as typhoid fever and cholera may be re- 
 moved from water by proper filtration or they may be 
 killed by boiling the water. A number of years ago there 
 was a serious epidemic of cholera in Hamburg. People 
 living in Hamburg used water taken directly from the river 
 Elbe. The people living in Altona, a city practically con- 
 tinuous with Hamburg, used water taken from the same 
 river but filtered it before it was pumped into the city mains. 
 Many people in Hamburg died from the cholera while there 
 were very few cases of cholera in Altona. In very many 
 cases epidemics of typhoid fever have been traced directly 
 to impure drinking waters. 
 
 Solutions. If salt or sugar is mixed with water, a certain 
 amount is dissolved. Substances which are not volatile 
 lower the freezing point and raise the boiling point of water 
 
CHEMICAL ACTIVITY. IONIZATION 65 
 
 or .any other solvent in which they are dissolved. The 
 lowering of the freezing point is familiar in the freezing 
 mixtures of ice and salt, which are used in making ice 
 cream. 
 
 Many substances are much more easily soluble in hot 
 water than in cold water. If a concentrated solution of such 
 a substance is cooled without the presence of any of the solid 
 substance, a supersaturated solution may be obtained, but 
 the introduction of a crystal of the substance will cause 
 a part of it to separate from the solution exactly as crystals 
 of ice will cause a part of supercooled water to freeze. A 
 supersaturated solution is unstable and cannot exist when 
 the solution and crystals of the substance are in intimate 
 contact. A solution obtained by shaking water or some 
 other solvent with a solid which dissolves in it till no more 
 is taken up is said to be saturated. 
 
 Chemical Activity in Solutions. lonization. When a 
 solution of salt, NaCl, is added to a solution of silver nitrate, 
 AgNO 3 , a white, curdy precipitate of silver chloride, AgCl, 
 is formed. Similar reactions will take place between dis- 
 solved substances in many other cases. Such actions 
 occur for two reasons: first, because in solution the sub- 
 stances are brought into intimate contact; second, because in 
 solution such substances as sodium chloride and silver 
 nitrate separate more or less completely into ions, the 
 sodium chloride, NaCl, into sodium, Na + , and chloride, 
 Cl~, ions; the silver nitrate into silver, Ag + , and nitrate, 
 N0 3 ~, ions, and any positive ion in a solution may combine 
 with any negative ion in the same solution. 
 
 Electrolysis. In considering the decomposition of water 
 by electricity (p. 43) it was stated that under the influence 
 of the electrical current the hydrogen of the sulfuric acid, 
 H 2 SO 4 , moves through the solution in one direction while 
 the sulfur and oxygen move together in the opposite direc- 
 tion. This is because the sulfuric acid separates into 
 
 5 
 
66 
 
 PROPERTIES AND USES OF WATER 
 
 positive hydrogen ions, H + , and negative sulf ate ions, SO j = . 
 The positive ions are attracted toward the cathode, where 
 they are discharged and liberated as free hydrogen, H 2 . 
 The sulfate ions, SO 4 = , are attracted toward the anode, 
 but in contact with the anode there are also a few hydroxyl, 
 OH~, and perhaps oxygen, = , ions, which are more easily 
 discharged than the sulfate ions. The hydroxyl or oxygen 
 ions are discharged, therefore, instead of the sulfate ions 
 and in this way free oxygen is liberated. 
 
 40H = 2H 2 O + O 2 
 
 Migration of Ions. The light hydrogen ions pass through 
 the solution toward the cathode much more rapidly than 
 the sulfate ions pass toward the anode. This different 
 rate of transfer for different ions can be demonstrated 
 by using the apparatus illustrated in Fig. 26. The U-tube 
 of the apparatus is filled with a solution containing 4 
 milligram molecules (146 mg.) of hydrochloric acid, divided 
 equally between the two arms. of the tube. A platinum 
 cathode and silver anode are used. The hydrogen ions, 
 
 The solution on this side 
 
 contains : 
 At first, 
 73 mg. HC1 = 
 2 milligram atoms of H 
 2 milligram atoms of Cl 
 At end, 
 
 66.5 mg. HC1 = 
 1.82 milligram atoms 
 
 of H 
 1.82 milligram atoms 
 
 of Cl 
 1 milligram atom of H 
 
 is liberated 
 
 Amount of hydrogen 
 transferred across the 
 line CD = 0.82 mg. 
 atoms. 
 
 Cathode 
 
 4- 
 
 V 
 
 i 
 
 m 
 
 3 
 
 = 
 
 / 
 
 
 ' 
 
 V 
 
 \^j *- 
 
 D 
 
 FIG. 26. 
 
 The solutior on this side 
 
 contains : 
 At first, 
 73 mg. HC1 = 
 2 milligram atoms of H 
 2 milligram atoms of Cl 
 At end, 
 43 mg. HC1 = 
 1.18 milligram atoms 
 
 of H 
 1.18 milligram atoms 
 
 of Cl 
 1 milligram atom of Cl 
 
 combines with anode 
 Amount ' of chlorine 
 transferred across the 
 line CD = 0.18 mg. 
 atoms. 
 
MIGRATION OF IONS 67 
 
 H + , discharged at the cathode are liberated as free hydrogen, 
 H 2 , while the chloride ions, Cl~, liberated at the anode 
 combine with the silver to form silvei chloride. 
 
 At the beginning of the experiment there is the same 
 amount of hydrochloric acid in each arm of the tube, 2 
 milligram molecules. If the current is passed through the 
 solution till 1 milligram atom of hydrogen has been 
 liberated at the cathode and 1 milligram atom of chlorine 
 has combined with the anode, it is evident that if the 
 hydrogen ions move through the solution in one direction 
 at the same rate that the chloride ions move the other 
 way there would be 1J^ milligram molecules of hydro- 
 chloric acid in each arm of the tube. Instead of this an 
 analysis of the solution in the two arms will show that the 
 cathode arm contains 1.82 milligram molecules of hydro- 
 chloric acid while the anode arm contains only 1.18 milli- 
 gram molecules. It is evident from this that 0.82 of the 
 milligram atom of hydrogen liberated at the cathode have 
 been replaced by hydrogen transferred from the anode side 
 while only 0.18 of the chlorine removed at the anode have 
 been replaced by chlorine transferred from the cathode 
 side. It follows that the hydrogen ion migrates toward 
 the cathode more than four times as fast as the chloride 
 ion migrates toward the anode. 
 
 The independent migration of the ions in a solution of 
 an electrolyte is one of the strong reasons for believing that 
 electrolytes actually separate into ions in a solution. 
 
 Mineral Waters. Rivers flowing into the ocean carry 
 small amounts of salt and other soluble compounds. Dur- 
 ing millions of years water has evaporated from the surface 
 of the ocean and the vapor has been carried back over the 
 land where it condenses and falls as rain, carrying a new 
 portion of mineral matters to the rivers and the ocean. 
 Some of the substances carried to the ocean are only slightly 
 soluble and these cannot accumulate to any large extent, but 
 
68 PROPERTIES AND USES OF WATER 
 
 salt and some other compounds are so easily soluble that the 
 ocean wa.ter is very far from a saturated solution of them. 
 In the Dead Sea, the Caspian Sea, Salt Lake and other 
 inland bodies of water with no outlet to the ocean the accu- 
 mulation of soluble salts has gone much further; so far, in- 
 deed, that even very soluble salts may be deposited from 
 them. During the long periods of geologic time vast beds 
 of such salts have been deposited and this has given rise 
 to the beds of rock salt found in New York, Kansas, South- 
 eastern Germany and elsewhere. In some cases potash 
 salts and other salts have been deposited and this is espe- 
 cially true in Stassfurt, Prussia, and in Alsace. For many 
 years the supply of potassium compounds for the world 
 came almost exclusively from Germany. 
 
 The waters of Searle's Lake in California and of some lakes 
 in Nebraska contain potassium salts, and during the great 
 war the recovery of compounds of potassium from these 
 sources has been rapidly developed. The prices charged 
 for these salts are, however, much higher than the prices 
 charged for potassium salts from Germany before the war. 
 
 Water which comes in contact with deposits of salt will, 
 of course, dissolve the soluble salts and in this way the 
 brines from which salt is made are formed at Syracuse, 
 New York; Midland, Michigan; and elsewhere. 
 
 Water containing other salts in solution may have valu- 
 able medicinal properties and such waters are sold as mineral 
 waters. Artificial waters similar in composition are often 
 made and sold. Some of these, as Appolinaris water, are 
 charged with carbon dioxide, to give them their effervescent 
 qualities. 
 
 Hydrates. Efflorescence, Deliquescence. Many salts 
 which are separated from solution in crystalline form by 
 evaporation combine with definite quantities of water. 
 Crystallized salts of this kind are called hydrates. Thus 
 the hydrate of copper sulfate has the composition CuSO 4 .- 
 
EFFLORESCENCE. DELIQUESCENCE 69 
 
 5H 2 0. The hydrate of sodium sulfate is Na 2 S0 4 .10H 2 O. 
 One of the hydrates of calcium chloride is CaCl 2 .6H 2 O. 
 
 Sodium sulfate, Na 2 S0 4 .10H 2 0, holds its water very 
 loosely and if exposed to ordinary air, which is not saturated 
 with moisture, the water of the salt will escape and the 
 partially dried salt will fall to a powder. This is called 
 efflorescence. The salts are said to be efflorescent. Other 
 salts, such as either anhydrous calcium chloride, CaCl 2 , or 
 the hydrate, CaCl 2 .6H 2 O, not only do not lose water in 
 ordinary air but they will even take up more water and 
 finally dissolve in the water absorbed, unless the air is very 
 dry indeed. Salts of this type are called deliquescent. 
 
 SUMMARY 
 
 The latent heat of fusion of ice is 79.63 calories. 
 
 The latent heat of evaporation of water is 536.6 calories. 
 
 Water has its maximum density at 4. 
 
 When water and its vapor are in contact the vapor has a 
 definite pressure which depends only on the temperature. 
 The temperature at which this vapor pressure is equal to 
 the pressure of the air is the boiling point. 
 
 Equilibrium between a liquid and its vapor is reached 
 when the number of molecules escaping from the liquid to 
 the vapor in a given time equals the number of molecules 
 returning to the liquid from the vapor. 
 
 Germs of such diseases as typhoid fever and cholera may 
 be carried by water. These may be removed by proper 
 filtration or killed by boiling the water or by sterilization. 
 
 A saturated solution contains all of a substance which the 
 liquid can retain while in contact with the solid phase. 
 Supersaturated solutions may be prepared in the absence of 
 the solid phase. 
 
 Substances in solution often react more readily than 
 others, partly because they are' in intimate contact, partly 
 
70 PROPERTIES AND USES OF WATER 
 
 because many substances separate into ions in solutions, 
 and the ions unite to form different compounds. 
 
 In electrolysis the positive ions, cations, move toward the 
 cathode, the negative ions, anions, toward the anode. The 
 rate of migration is different for different ions. 
 
 Mineral waters are found in many places and are some- 
 times valuable for their medicinal properties. 
 
 Many salts and sometimes other compounds combine 
 with water to form hydrates. Some of these are efflo- 
 rescent, others deliquescent. 
 
 EXERCISES 
 
 1. What per cent of water is contained in crystallized copper 
 sulfate? In crystallized sodium sulfate? In crystallized calcium 
 chloride? 
 
 2. How many grams of copper will be required to give a pound 
 of crystallized copper sulfate? Of anhydrous copper sulfate? 
 
 3. What method is used in your community to secure a water 
 supply free from disease germs? 
 
 4. The following methods of purifying or sterilizing water have 
 been used: 
 
 (a) Slow sand nitration. 
 
 (6) Addition of aluminium sulfate (so-called nitration "alum ") 
 followed by rapid, mechanical nitration. 
 
 (c) Aluminium sulfate and "lime" (calcium hydroxide) 
 
 followed by settling basins and filtration. 
 
 (d) Ozone. 
 
 (e) Ultra-violet light. 
 (/) Bleaching powder. 
 (g) Liquid chlorine. 
 
 What are some of the advantages and disadvantages of each of 
 these methods? 
 
 5. The amount of aluminium sulfate used in filtration plants is 
 from 0.3 to 2 grains per gallon. How many gallons can be puri- 
 fied by the use of a pound? One pound is 7000 grains and one 
 gallon is 8H (8.3389) pounds. 
 
EXERCISES. WATER 71 
 
 6. From 0.2 to 0.5 parts per million of liquid chlorine are used 
 to sterilize water. How many gallons can be sterilized by a pound? 
 How many gallons can be sterilized by a pound of bleaching powder 
 containing 33^ per cent of available chlorine? 
 
 7. What has been the effect of the drainage canal on the typhoid 
 death rate in Chicago? 
 
 8. How many deaths were there, per thousand soldiers, in the 
 cantonments of the Spanish-American War? In the Russo- 
 Japanese War? In the Great War? What are the causes of these 
 differences? 
 
 9. By what means other than contaminated water is typhoid 
 fever communicated? What means were used to prevent typhoid 
 fever in the armies of the Great War? 
 
CHAPTER VII 
 gODIUM, ACIDS, BASES, SALTS 
 
 Salt. Common salt is a substance composed of sodium, 
 Na, and chlorine, Cl, and has the composition represented 
 by the formula NaCl. The chemical name is sodium 
 chloride. Salt is very widely diffused. It is a necessary 
 constituent of the food which we eat and is found in small 
 amounts in practically all natural waters which have come 
 In contact with the earth. The salty taste of sea water is 
 cniefly due to sodium chloride, though sea water also con- 
 tains a number of other chlorides. Thick beds of rock salt 
 are found in Germany, Louisiana, Kansas and many other 
 places. Strong brines, which are nearly saturated solutions 
 of sodium chloride, are also obtained by means of artesian 
 wells, 'and salt is manufactured by evaporating these brines 
 at Syracuse, New York, in Michigan, and at many other 
 places in the world. 
 
 Electrolysis of Fused Sodium Chloride. If salt is melted 
 in a crucible and an electric current is passed through it 
 between carbon electrodes, chlorine gas will be evolved at 
 the anode ancj sodium will be liberated at the cathode. 
 
 If a current is passed through a solution of sodium 
 chloride in water, the chloride ions, Cl~~, will be carried 
 toward the anoole by the current and the sodium ions, Na + , 
 will be carried toward the cathode. In contact with the 
 anode the chloride ions will be discharged and chlorine 
 gas will be evolved. In any aqueous solution there are 
 present, besides the ions of the dissolved salt, hydrogen 
 ions H + , and hydroxide ions, OH~, which come from the 
 /( -- U 72 
 
SODIUM. ACIDS 73 
 
 ionization of the water, H 2 O (or HOH). Hydrogen ions 
 require much less electrical energy than sodium ions for 
 their discharge. At the cathode, therefore, the hydrogen 
 ions are discharged and hydrogen gas is liberated, while 
 the hydroxide ions, OH~, and sodium ions, Na + , remain in 
 solution, forming a solution of sodium hydroxide, NaOH. 
 This method is now extensively used for the manufacture 
 of sodium hydroxide and also of chlorine. 
 
 Sodium is a soft metal which looks very much like silver 
 when it is first cut. The surface tarnishes almost instantly 
 in the air, however, because it decomposes the moisture in 
 the air, liberating hydrogen and forming sodium hydroxide : 
 
 2Na + 2HOH = 2NaOH + H 2 
 
 The same reaction occurs when a small piece of sodium 
 is thrown on water. The metal is lighter than water and 
 floats on the surface, melting to a globule from the heat of 
 the reaction. This globule glides rapidly over the surface 
 of the water as hydrogen is evolved at the surface of contact. 
 The sodium hydroxide, NaOH, dissolves in the water and 
 gives to it an alkaline reaction, that is, it will turn red 
 litmus blue, and it will neutralize acids (p. 75). 
 
 Acids. In a previous chapter several acids, such as 
 hydrochloric acid, HC1, sulfuric acid, H 2 SO 4 , and acetic 
 acid, HC 2 H 3 O 2 , were spoken of. It was pointed out that 
 the element common to all acids is hydrogen and that this 
 hydrogen may be replaced by metals, as in the action of 
 sulfuric acid on zinc, which gives zinc sulfate, ZnSO 4 , and 
 hydrogen. 
 
 Another very important characteristic of acids is that 
 they separate, or ionize, in solutions in water, giving hydro- 
 gen ions which have a positive charge and some negative 
 atom or group of atoms. Thus hydrochloric acid gives 
 hydrogen ions, H+, and chloride ions, Cl~; sulfuric acid 
 gives hydrogen ions, H+, H+, and sulfate ions, S04 = ; acetic 
 
74 SODIUM, ACIDS, BASES, SALTS 
 
 acid, HC 2 H 3 O 2 , gives hydrogen ions, H + , and acetate ions, 
 C 2 H 3 O 2 ~. If a current of electricity is passed through a 
 solution of an acid, the hydrogen ions are carried by the 
 current toward the cathode, or negative electrode, while 
 the negative ions are carried toward the anode, or positive 
 electrode. 
 
 Bases. Sodium hydroxide, NaOH, belongs to a class of 
 compounds called bases. These ionize in solution, giving 
 hydroxide ions, OH~, carrying a negative charge. Thus 
 sodium hydroxide gives hydroxide ions, OH~, and sodium 
 ions, Na + ; calcium hydroxide, Ca(OH) 2 , gives hydroxide 
 ions, OH~, OH~, and calcium ions, Ca ++ , ammonium hy- 
 droxide, NH 4 OH, gives hydroxide ions, OH~, and ammonium 
 ions, NH 4 +. 
 
 From the point of view just presented the most essential 
 characteristic of an acid is that it gives hydrogen ions in 
 solution, and the most characteristic property of a base is 
 that its solution contains hydroxide ions. 
 
 lonization of Water. Neutralization. Salts. Pure water, 
 H 2 0, or HOH, separates to a very slight extent into hydro- 
 gen ions, H+, and hydroxide ions, OH~. This may be rep- 
 resented by the equilibrium equation: 
 
 HOH *=> H+ + OH- 
 
 In this reaction the equilibrium is very far to the left, 
 i.e., there are always a very large number of molecules of 
 water in comparison with the number of hydrogen and 
 hydroxide ions (in pure water approximately 500,000,- 
 000 : 1) If a solution of an acid, containing hydrogen 
 ions, H + , is added to a solution of a base, containing hydrox- 
 ide ions, OH~, the two kinds of ions must unite with each 
 other until either the hydrogen or the hydroxide ions, or 
 both, nearly disappear. If the acid and base are used in 
 equivalent amounts, the resulting solution must contain 
 equal numbers of hydrogen and hydroxide ions and the 
 
NEUTRALIZATION 75 
 
 number of each of these ions must be very small. Such 
 a solution is said to be neutral. A base is said to neutralize 
 an acid because each destroys the characteristic property 
 of the other by the process just described. The process of 
 neutralization gives rise to reactions which may be illus- 
 trated by such equations as the following: 
 
 NaOH + HC1 = NaCl + H 2 O 
 
 Sodium 
 chloride 
 
 2NaOH + H 2 SO 4 = Na 2 S0 4 + 2H 2 
 
 Sodium 
 sulfate 
 
 Ca(OH) 2 + 2HC 2 H 3 2 = Ca(C 2 H 3 O 2 ) 2 + 2H 2 O 
 
 Calcium 
 acetate 
 
 2A1(OH) 3 + 3H 2 SO 4 = A1 2 (SO 4 ) 3 + 3H 2 O 
 
 Aluminium 
 sulfate 
 
 It is to be noticed that in each equation the number of 
 hydrogen atoms in the acid must be the same as the number 
 of hydroxide groups in the base and that this fixes the for- 
 mula of the salt. 
 
 The compound formed by the union of the positive ion 
 of a base and the negative ion of an acid is called a salt. 
 A salt is also frequently defined as a compound formed by 
 the replacement of the hydrogen of an acid by a metal. 
 
 Exercise. Write the sixteen equations representing the 
 reactions which may occur between the following acids 
 and bases: 
 
 Hydrochloric acid, HC1 Sodium hydroxide, NaOH 
 
 Nitric acid, HN0 3 Ammonium hydroxide, NH 4 OH 
 
 Sulfuric acid, H 2 S0 4 Calcium hydroxide, Ca(OH) 2 
 
 Phosphoric acid, H 3 PO 4 Ferric hydroxide, Fe(OH) 3 
 
 Indicators. A number of natural and artificial dyes 
 exhibit a different color in an acid from that shown 'in an 
 
76 SODIUM, ACIDS, BASES, SALTS 
 
 alkaline solution. Thus in an acid solution, that is, in a 
 solution which contains more than a very few hydrogen 
 ions, litmus is red, while in an alkaline solution it becomes 
 blue. Phenol phthalein, on the other hand, is red in an 
 alkaline solution but colorless in an acid solution. The 
 change in color is due to a change in composition of the dye, 
 but the nature of the change cannot be given here. Dyes 
 of this character are called indicators, and they are much 
 used to determine the " end-point'* when an acid is neu- 
 tralized by a base. Paper which has been dipped in a so- 
 lution of litmus and dried is used to determine whether a 
 given solution is acid or alkaline. 
 
 Dibasic Acids. Either one or both of the hydrogen atoms 
 of sulfuric acid may be replaced by a metal, giving acid and 
 normal salts, as acid sodium sulfate, NaHS0 4 , and normal 
 sodium sulfate, Na 2 SO 4 . Acids having this property are 
 called dibasic. An acid like phosphoric acid, H 3 PO 4 , 
 which forms three salts with sodium, NaH 2 PO 4 , Na 2 HP0 4 
 and Na 3 P0 4 , is called tribasic. The basicity depends, 
 however, not on the number of hydrogen atoms in one 
 molecule of an acid, but on the number of replaceable 
 hydrogen atoms. Thus acetic acid, C 2 H 4 O 2 , is monobasic 
 because only one of its hydrogen atoms can be replaced; 
 and phosphorous acid, H 3 PO 3 , is dibasic because only two 
 of the hydrogen atoms can be replaced. 
 
 Normal, Neutral and Acid Salts. When a solution of a 
 base is mixed with an equivalent amount of a solution of 
 an acid the hydrogen, H + , and hydroxide, OH~~, ions unite to 
 form water, H 2 0. The metallic ions of the base, such as 
 Na + , from NaOH, and the non-metallic ions of the acid, 
 such as Cl~, from hydrochloric acid, HC1, or SO 4 = from sul- 
 furic acid, H 2 SO 4 , partly remain in solution as ions and 
 partly unite to form salts. In the cases referred to the 
 salts are sodium chloride, NaCl, and sodium sulfate, Na 2 S0 4 . 
 If the solution is evaporated to dryness the ions will all 
 
SALTS OF WEAK ACIDS 77 
 
 unite to form the salt. In other words, in the reversible 
 reactions represented by the equations: 
 
 NaCl => Na+ + Gl- 
 and: 
 
 Na 2 SO 4 *=> Na+ + Na+ + SO 4 = 
 
 the separation into ions is favored by the addition of more 
 and more water, and electrolytes, such as the salts named, 
 and strong acids, like hydrochloric or sulfuric acid, are 
 almost completely ionized in dilute solutions. On the other 
 hand, the concentration of the solution by the removal of 
 water causes the ionization to decrease and in the dry salt 
 very few, if any, ions remain. 
 
 Dibasic acids form acid salts such as acid sodium sulfate, 
 NaHSO 4 , and normal salts, as sodium sulfate, Na 2 S0 4 . 
 Normal salts of strong acids, that is, of acids which are 
 largely ionized in dilute solution, with strong bases, are 
 neutral in their reaction toward litmus and other indicators. 
 This means, of course, that the numbers of hydrogen, 
 H+, and hydroxide, OH , ions in solutions of such salts 
 are equal. 
 
 Normal salts of weak acids, such as carbonic acid, H 2 CO 3 , 
 on the other hand, are frequently 'alkaline toward litmus. 
 Sodium carbonate, Na 2 CO 3 , ionizes in solution to sodium 
 ions, Na+, Na + , and carbonate ions, C0 3 = . The carbonate 
 ions, however, combine with the hydrogen ions of the water 
 to form bicarbonate ions, HC0 3 ~: 
 
 CO 3 = + H+ <=> HCO-r 
 
 Since the bicarbonate ions have only a slight tendency to 
 separate into hydrogen ions and carbonate, CO 3 = , ions, be- 
 cause carbonic acid is a very weak acid, this reaction removes 
 a considerable number of hydrogen ions from the solution. 
 These will be replaced by the further ionization of water: 
 
 HOH<=H++ OH- 
 
78 SODIUM, ACIDS, BASES, SALTS 
 
 The process leaves an excess of hydroxide ions, OH~, in 
 the solution and the reaction is, therefore, alkaline. 
 
 Because of these properties of salts of weak acids those 
 salts in which all of the replaceable hydrogen of an acid has 
 been displaced by a metal are called normal salts, rather 
 than neutral salts. 
 
 Some confusion is liable to arise because salts in which a 
 part only of the hydrogen of an acid has been replaced are 
 called add salts, although such salts may be neutral or even 
 alkaline in reaction. Thus NaHCO 3 is called, frequently, 
 acid sodium carbonate, though it is neutral in reaction to 
 litmus. This use of the term add salt is based on the older 
 definition of an acid as a compound containing replaceable 
 hydrogen, and is not consistent with the modern definition 
 of an acid as a compound whose solution contains more hydro- 
 gen than hydroxide ions. 
 
 SUMMARY 
 
 Common salt or sodium chloride is our most common 
 easily soluble salt. 
 
 The electrolysis of melted salt, NaCl, gives metallic 
 sodium and chlorine. 
 
 The electrolysis of a solution of salt gives sodium hydrox- 
 ide and hydrogen at the cathode and chlorine at the anode. 
 
 Acids are hydrogen compounds which ionize to positive 
 hydrogen and negative ions. They are also defined as 
 hydrogen compounds in which the hydrogen may be re- 
 placed by a metal. 
 
 Bases are hydroxyl (OH) compounds which ionize to 
 negative hydroxide ions and some positive ion, which is 
 usually metallic. 
 
 Neutralization consists in the union of the hydrogen ions 
 of an acid with the hydroxide ions of a base, forming water, 
 while the negative ion of the acid and the positive ion of the 
 
EXERCISES. ACIDS, BASES 79 
 
 base either remain in solution uncombined or unite to form a 
 salt. 
 
 In a neutral solution the numbers of hydrogen and hydrox- 
 ide ions are equal. 
 
 An indicator exhibits one color in the presence of an ex- 
 cess of hydrogen ions and another color in the presence of 
 an excess of hydroxide ions. 
 
 Normal salts are salts in which all of the replaceable 
 hydrogen of an acid has been replaced by a metal or some 
 other metallic group. Normal salts may be neutral, alkaline 
 or acid in reaction. 
 
 Acid salts, in the older use of the term, are salts in which 
 only a part of the replaceable hydrogen has been replaced 
 by a metal. 
 
 EXERCISES 
 
 1. How many grams of hydrochloric acid must there be in a 
 liter of a solution of the acid which will exactly neutralize a liter 
 of a solution containing 40 grams of sodium hydroxide? 
 
 2. How many grams of sulfuric acid in one liter will give a 
 solution exactly equivalent to the sodium hydroxide solution 
 just mentioned? How many grams of nitric acid? Of acetic 
 acid? 
 
 3. How much lime, CaO, will be neutralized by one liter of any 
 one of the acid solutions mentioned above? 
 
 4. What-is the difference between the method of preparing salts 
 given in this chapter and that given in Chapter IV? 
 
 5. Write the equation for the reaction between acetic acid and 
 calcium. For the reaction between sulfuric acid and sodium. 
 
 6. Why does a globule of melted sodium move rapidly over the 
 surface of water? 
 
CHAPTER VIII 
 
 HYDROCHLORIC ACID. CHLORINE. OXYGEN 
 ACIDS OF CHLORINE 
 
 Preparation of Hydrochloric Acid. The addition of con- 
 centrated sulfuric acid, H 2 S0 4 , to salt, NaCl, causes the 
 liberation of hydrochloric acid as a gas : 
 
 NaCl + H 2 S0 4 ?=* NaHSO 4 + HG1 
 
 Sodium Sulfuric Acid Hydrochloric 
 
 chloride acid sodium acid 
 
 sulfate 
 
 The compound NaHSO4 is called acid sodium sulfate be- 
 cause it still has acid properties. How might this be shown? 
 
 It is sometimes stated that hydrochloric acid is liberated 
 because sulfuric acid is a stronger acid than hydrochloric- 
 Such a statement overlooks the fact that the addition of 
 concentrated hydrochloric acid to a concentrated solution 
 of acid sodium sulfate causes a precipitation of salt, NaCl. 
 
 NaHSO 4 + HC1 <=> NaCl + H 2 SO 4 
 
 If sulfuric acid is the stronger acid in one case, hydro- 
 chloric acid is stronger in the other. It is clear, therefore, 
 that the direction of the reaction does not depend on the 
 relative strengths of the two acids. It depends, instead, 
 on the fact that the reaction is reversible and in the first 
 case it goes toward the liberation of hydrochloric acid 
 because that escapes as a gas, and in the second case it goes 
 toward the formation of salt because the salt is removed 
 from the sphere of action as a precipitate. 
 
 Properties of Hydrochloric Acid. Hydrochloric acid is a 
 colorless gas, about one-fourth heavier than air. It has a 
 
 80 
 
ACTION ON METALS 81 
 
 very pungent, irritating odor, even when diluted with a 
 large amount of air. It dissolves in water and at the freez- 
 ing point one volume of water will take up 500 volumes 
 of the gas. The solution contains forty-five per cent of 
 hydrochloric acid. 
 
 This concentrated solution gives off some of the gas at 
 ordinary temperatures. If the solution is boiled, hydro- 
 chloric acid escapes chiefly at first and the solution becomes 
 less and less concentrated, till there finally remains a 
 solution which contains only twenty per cent of hydro- 
 chloric acid and which boils at about 110. On the other 
 hand, a dilute solution containing less than twenty per 
 cent of the acid allows water to escape at first, when 
 boiled, and the boiling point gradually rises to 110, and 
 the concentration of the acid remaining increases to twenty 
 per cent. This acid, which boils constantly at 110, may 
 be obtained, as described, from either a more concentrated 
 or from a more dilute acid. 
 
 When hydrochloric acid escapes into the air a cloud is 
 formed unless the air is very dry. This is because the acid 
 forms with the water of the air a mixture or compound 
 having a higher boiling point than the boiling point of 
 pure water. This mixture will condense to minute drops 
 of a concentrated solution of hydrochloric acid at a tem- 
 perature at which the water contained in the volume of air 
 occupied by the cloud would all remain as an invisible 
 vapor. 
 
 Action of Hydrochloric Acid on Metals. The solution 
 of hydrochloric acid in water acts on many of the metals, 
 forming chlorides of the metals while the hydrogen of the 
 acid is liberated as a gas. In this way sodium gives sodium 
 chloride, NaCl; zinc gives zinc chloride, ZnCl 2 ; iron gives 
 ferrous chloride, FeCl 2 ; aluminium gives aluminium 
 chloride, A1C1 3 , and tin gives stannous chloride, SnCl 2 . 
 
82 HYDROCHLORIC ACID 
 
 The equations for these reactions should be written by 
 the student. 
 
 Action of Hydrochloric Acid on Hydroxides and Oxides 
 of the Metals. The formation of salts by the Action of 
 acids on bases was discussed in the last chapter and should 
 be recalled here. 
 
 In a somewhat similar manner many oxides of the metals 
 dissolve in a solution of hydrochloric acid, giving chlorides 
 of the metals and water. Thus zinc oxide gives zinc 
 chloride : 
 
 ZnO + 2HC1 = ZnCl 2 + H 2 O 
 
 Zinc Zinc 
 
 oxide chloride 
 
 It is to be noticed that in reactions of this character two 
 molecules of hydrochloric acid are required for each atom 
 of oxygen in the oxide. In a similar manner ferric oxide, 
 Fe 2 O 3 , cuprous oxide, Cu 2 O, cupric oxide, CuO, mercuric 
 oxide, HgO, and bismuth oxide, Bi 2 O 3 , give corresponding 
 chlorides. The magnetic oxide of iron, Fe 3 O 4 , gives a mix- 
 ture of ferrous chloride, FeCl 2 , and ferric chloride, FeCl 3 . 
 The student should write the equations for all of these 
 reactions. 
 
 Action of Oxidizing Agents on Hydrochloric Acid. By 
 heating a mixture of hydrochloric acid with air or oxygen a 
 part of the hydrogen of the acid is oxidized to water and 
 chlorine is liberated. The reaction is reversible and may be 
 represented thus: 
 
 4HC1 + O 2 =F 2H 2 O + 2C1 2 
 
 By using copper chloride as a catalytic agent (p. 9) the 
 reaction has sometimes been used for the commercial 
 preparation of chlorine, but as a technical process this 
 has not been able to compete with other methods of obtain- 
 ing the element. 
 
 The oxidation of hydrochloric acid may also be effected 
 
PROPERTIES OF CHLORINE 83 
 
 by a number of oxidizing agents at ordinary temperatures, 
 or by gently warming a concentrated solution of hydrochloric 
 acid with the agent. The substances most often used in 
 the laboratory for this purpose are manganese dioxide, 
 MnO 2 , potassium permanganate, KMnO 4 , and potassium 
 dichromate, K 2 Cr 2 O 7 . In the first case manganese chloride, 
 MnCl 2 , is formed, in the second case manganese chloride 
 and potassium chloride: 
 
 Mn0 2 + 4HC1 = MnCl 2 + 2H 2 O + C1 2 
 2KMnO 4 + 16HC1 = 2MnCl 2 + 2KC1 -f 8H 2 O + 5C1 2 
 
 It should be noticed that in each case enough molecules 
 of hydrochloric acid must be used to furnish hydrogen to 
 combine with all of the oxygen of the oxidizing agent. 
 
 Properties of Chlorine. Chlorine forms compounds with 
 nearly all of the chemical elements, and in many cases the 
 action of chlorine upon an element takes place at a much 
 lower temperature than that at which the same element 
 will combine with oxygen. Finely powdered antimony 
 will burst into flame when sprinkled through the gas, and 
 false gold leaf, an alloy of copper and zinc, will also flash 
 up when the gas is passed over it. The antimony forms 
 antimony trichloride, SbCl 3 , or antimony pentachloride, 
 SbCl 5 , according as the antimony or chlorine is in excess. 
 Phosphorus forms similar compounds, phosphorus tri- 
 chloride, PC1 3 , and phosphorus pentachloride, PC1 5 . The 
 copper and zinc of the false gold leaf form cuprous chloride, 
 Cu 2 Cl 2 , or cupric chloride, CuCl 2 , and zinc chloride, ZnCl 2 . 
 
 Chlorine gas is very poisonous and great care should be 
 taken to avoid breathing it. When it has been inhaled 
 in small quantities the best antidote for immediate use 
 is vapor of alcohol. The gas masks used in warfare contain 
 soda-lime and sodium thiosulfate. For the absorption 
 of other poisonous gases charcoal prepared from cocoanut 
 shells is used. 
 
84 HYDROCHLORIC ACID 
 
 Bleaching. Many vegetable and artificial colored sub- 
 stances are bleached by moist chlorine gas or by a solution 
 of chlorine in water. The brown colors which are charac- 
 teristic of unbleached cotton and linen are also destroyed 
 in this way. These colors are sensitive to sunlight, and be- 
 fore the method of bleaching with chlorine was discovered 
 it was customary to spread cotton or linen cloth on the 
 grass and expose it to the action of the sunlight for many 
 days or weeks to secure the pure white color which was 
 desired. The same effect is now obtained very quickly 
 by the use of a dilute solution of chlorine. Dry chlorine 
 acts very slowly or not at all on the colored calico. Because 
 of this fact it is believed that the chlorine does not act 
 directly on the moist colored substances but that it acts 
 at first on the water, forming hypochlorous acid, HC10, 
 and hydrochloric acid: 
 
 C1 2 + HOH = HC1 + HOC1 
 
 Hydrochloric Hypochlorous 
 acid acid 
 
 The hypochlorous acid gives its oxygen readily to other 
 substances and seems to be the actual bleaching agent. 
 It oxidizes the colored compounds to other compounds 
 which are colorless. 
 
 Names of Oxygen Acids of Chlorine and of Their Salts. 
 Chlorine combines with hydrogen and oxygen to form four 
 different acids. Each of these contains one atom of hydro- 
 gen and one atom of chlorine in its molecule and, with these, 
 one, two, three or four atoms of oxygen. The names of 
 these acids are: 
 
 Hypochlorous acid, HC10 
 
 Chlorous acid, HC102 
 
 Chloric acid, HC10 3 
 
 Perchloric acid, HC1O 4 
 
 In these names the endings ous and ic are used as they 
 are used in naming oxides or chlorides (p. 48), the com- 
 
NOMENCLATURE. HYPOCHLOROUS ACID 85 
 
 pound with the ending ous containing less, and the one 
 having the ending ic containing more oxygen. The pre- 
 fix hypo means under or less, and the prefix per means above 
 or beyond. Hypochlorous acid contains less oxygen than 
 chlorous acid; perchloric acid contains more oxygen than 
 chloric acid. 
 
 Corresponding to these acids are many salts in which 
 the hydrogen of the acids is replaced by metals. Thus 
 the sodium salts are: 
 
 Sodium hypochlorite, NaCIO 
 
 Sodium chlorite, NaClO 2 
 
 Sodium chlorate, NaClO 3 
 
 Sodium perchlorate, NaC10 4 
 
 In these names the ous of the acids is changed to ite for the 
 salts and the ic of the acids is changed to ate for the salts. 
 
 In accordance with these principles name the following 
 compounds : 
 
 HBrO, HBr0 3 , KBrO, KBr0 3 , KIO 3 , KIO 4 , H 2 SO 3 , 
 
 H 2 S0 4 , Na 2 S0 3 , Na 2 S0 4 , H 3 P0 3 , H 3 P0 4 , K 3 P0 4 . 
 
 Hypochlorous Acid. Effect of Removing One of the 
 Products of a Reversible Reaction. When chlorine is dis- 
 solved in water a small part of it reacts with the water in 
 accordance with the reversible reaction expressed by the 
 equation : 
 
 C1-C1 + H-OH 1 <=> H-C1 + HO-C1 (or HC10) 
 
 Hydrochloric Hypochlorous 
 
 acid acid 
 
 In this reaction the equilibrium, which is always reached 
 sooner or later in a reversible reaction, is far to the left. In 
 other words, there will be a much larger proportion of chlo- 
 rine (C1 2 ) in the solution than there will be of hydrochloric 
 
 1 The formulas of chlorine, Cl2, and water, HaO, are written in this 
 manner to call attention to the way in which each substance divides in 
 reacting with the other. 
 
86 HYDROCHLORIC ACID 
 
 and hypochlorous acids, because the reaction between the 
 hydrochloric acid and hypochlorous acid giving chlorine 
 and water goes very much faster than that between chlorine 
 and water giving hypochlorous acid and hydrochloric acid. 
 If a solution of potassium hydroxide or of any other 
 strong alkali is added to such a solution, the base reacts 
 with the acids in accordance with the usual interaction of 
 acids and bases: 
 
 HC1 + KOH * KC1 + HOH 
 HC10 + KOH <= KC10 + HOH 
 
 The reactions shown by these equations are also reversi- 
 ble with the equilibrium very far to the right, so that very 
 little of the acids can remain in a solution containing a base. 
 As soon, however, as the hydrochloric and hypochlorous 
 acids are neutralized by the base the chlorine and water 
 will react to form more of the acids, and all of the chlorine 
 will be very quickly converted into a mixture of potassium 
 chloride, KC1, and potassium hypochlorite, KC1O. As 
 these substances are the first tangible products of the 
 interaction of chlorine and potassium hydroxide we may 
 add the three equations together, omitting compounds 
 which occur on opposite sides : 
 
 C1 2 + HOH = HC1 + HC10 
 HC1 + KOH = KC1 + HOH 
 HC10 + KQH = KC1Q + HOH 
 01, + 2KOH = KC1 + KC10 + H 2 O 
 
 Bleaching Powder or Chloride of Lime. In calcium 
 hydroxide, Ca(OH) 2 , or slaked lime, there are two hydroxide 
 groups, OH, in the same molecule. When chlorine reacts 
 
 Cl 
 
 with calcium hydroxide, therefore, a compound, Ca<^ , 
 
 CIO 
 
BLEACHING POWDER. CHLORIC ACID 8? 
 
 is formed which is partly chloride and partly hypochlorite: 
 
 OH Cl 
 
 Ca/ + C1 2 = Ca/ + H^O 
 OH CIO 
 
 Bleaching 
 powder 
 
 Bleaching powder is intermediate between calcium 
 Cl CIO 
 
 chloride, Ca\^ , and calcium hypochlorite, Ca<^ : 
 Cl CIO 
 
 Bleaching powder or chloride of lime has proved to be the 
 most convenient method of transporting chlorine for most 
 of the technical uses to which it is applied. It is a solid 
 which can be readily carried in bulk in ordinary receptacles 
 and when chlorine is wanted for any purpose the addition 
 of an acid to the bleaching powder causes the reversal of 
 the reactions which have been given above and liberates, 
 chlorine : 
 
 Cl 
 
 Ca/ + H 2 SO 4 = CaSO 4 4 HC1 -f HC1O 
 
 OC1 
 
 HC1O -r HC1 = C1 2 + HOH 
 
 Sulfuric acid is usually chosen to liberate the chlorine? 
 because it is the cheapest of the acids. 
 
 Chloric Acid. Potassium Chlorate. It will be remem-* 
 bered that potassium chlorate, KClOs, decomposes easily, 
 when heated, into potassium chloride, KC1, and oxygen. 
 This decomposition is used as the simplest laboratory 
 method of preparing oxygen. 
 
 The hypochlorites all decompose, even more easily, 
 chlorides and oxygen: 
 
 2KC1O = 2KC1 + 0, 
 
 Potassium Potassium 
 
 hypochlorite chloride 
 
88 HYDROCHLORIC ACID 
 
 The hypochlorites may also take up oxygen and pass into 
 the more stable chlorate: 
 
 KC10 + 2O = KC1O 3 
 
 Potassium 
 chlorate 
 
 In a warm solution, as, for instance, when chlorine gas 
 is passed into a concentrated solution of potassium hydrox- 
 ide, the two reactions may occur simultaneously in accord- 
 ance with the equation: 
 
 2KC10 + KC10 = KC10 3 + 2KC1 
 
 Two molecules of potassium hypochlorite give their oxy- 
 gen to a third molecule: 
 
 KC1 
 KC1 
 
 O 
 
 KC10 
 
 If we add this equation to three times the equation given 
 above to represent the action of chlorine on a cold solution 
 of potassium hydroxide, canceling the potassium hypo- 
 chlorite, KC10, from both sides, we have the following: 
 
 6KOH + 3C1 2 = 3KC10 + 3KC1 + 3H 2 
 
 3KC1O + = KC1O 3 + 2KC1 
 
 6KOH + 3C1 2 = KClOa + 5KC1 + 3H 2 
 
 This last equation may be taken as representing what 
 occurs when chlorine gas is absorbed by a warm solution 
 of potassium hydroxide. 
 
 When a strong acid is added to a solution of potassium 
 chlorate some of the chloric acid is liberated in the solution : 
 
 2KC10 3 + H 2 SO 4 ^ 2HC10 3 + K 2 S0 4 
 
 If such a solution is evaporated, however, the chloric acid 
 decomposes into water, chlorine, chlorine dioxide, C1O 2 , 
 and other substances, and it has never been found possible 
 
POTASSIUM PERCHLORATE 89 
 
 to prepare chloric acid as a pure compound, free from water. 
 If concentrated sulfuric acid is added to potassium chlorate, 
 these decompositions take place at once, and chlorine dioxide 
 may cause the ignition of sugar or other organic matter 
 in contact with the mixture. Matches were once made 
 on the basis of these facts but they were soon displaced by 
 the more convenient phosphorus matches. 
 
 Potassium Perchlorate. The decomposition of potas- 
 sium chlorate by heat for the preparation of oxygen will 
 be remembered (p. 9). If potassium chlorate is heated 
 just above its melting point for some time, one portion 
 of the salt gives oxygen to another portion, converting it 
 into potassium perchlorate: 
 
 KC10 3 + 3KC10 3 = KC1 + 3KC1O 4 
 
 Potassium 
 perchlorate 
 
 One molecule of potassium chlorate gives enough oxygen 
 to convert three molecules of the chlorate to the perchlorate. 
 
 Potassium perchlorate is much more difficultly soluble 
 in water than potassium chlorate. It requires a higher 
 temperature for its decomposition into potassium chloride 
 and oxygen. It is noticeable that the acids of chlorine 
 containing oxygen and their salts become more and more 
 stable as the amount of oxygen increases. Potassium 
 chlorate is very much more stable than potassium hypo- 
 chlorite and the perchlorate is more stable than the chlo- 
 rate. The increase in stability is still more marked in the 
 case of the free acids. 
 
 Perchloric Acid, HC1O 4 , may be distilled under low 
 pressures without decomposition. 
 
 If a solution of perchloric acid is added to a solution of 
 potassium chloride or of some other soluble potassium 
 salt, a precipitate will be formed because potassium per- 
 chlorate is only very slightly soluble: 
 
 KC1 + HC1O 4 = KC10 4 + HC1 
 
90 HYDROCHLORIC ACID 
 
 This property of perchloric acid is used for the detection 
 and also for the quantitative determination of potassium. 
 
 SUMMARY 
 
 Hydrochloric acid is a gas which is prepared by the action 
 of sulfuric acid on salt. 
 
 Hydrochloric acid dissolves easily in water. A solution 
 containing 20 per cent of the acid boils constantly at 110. 
 A very concentrated solution gives chiefly hydrochloric acid 
 when boiled, while a dilute solution gives mainly water. 
 
 Hydrochloric acid gives with zinc, iron, tin and other 
 metals hydrogen and a chloride of the metal. 
 
 Hydrochloric acid gives with oxides or hydroxides of 
 metals water and a chloride of the metal. 
 
 Oxidizing agents oxidize the hydrogen of hydrochloric 
 acid, liberating chlorine. 
 
 Chlorine is an active element combining directly both 
 with metals and with non-metals. 
 
 Moist chlorine bleaches nearly all vegetable colors, de- 
 stroying them. 
 
 A logical nomenclature for acids and salts is developed 
 by use of the endings ous and ic, ite and ate t and the prefixes 
 hypo and per. 
 
 The primary action of chlorine with water gives hypo- 
 chlorous and hydrochloric acids. 
 
 Hypochlorites are prepared by the action of chlorine on 
 hydroxides, a chloride being formed at the same time. 
 
 Chlorine gives with calcium hydroxide a compound which 
 is both hypochlorite and chloride. 
 
 Potassium hypochlorite when warmed in solution gives 
 potassium chloride and potassium chlorate. 
 
 When potassium chlorate is heated to the melting point 
 it is partly changed to a mixture of potassium perchlorate 
 and potassium chloride. 
 
EXERCISES. CHLORINE 91 
 
 Potassium perchlorate is precipitated when a solution of 
 perchloric acid is added to a solution of a potassium salt. 
 
 EXERCISES 
 
 1. Write the equations for the following: 
 (a) Calcium chloride and sulfuric acid. 
 
 (6) Chlorine and a cold solution of sodium hydroxide. 
 
 (c) Chlorine and a warm solution of sodium hydroxide.. 
 
 (d) The effect of boiling a solution of bleaching powder. 
 
 (e) Potassium sulfate and perchloric acid. 
 (/) Sodium chlorate when heated. 
 
 2. If slaked lime is treated with chlorine and the water formed 
 does not escape, what per cent of chlorine will the product 
 contain? 
 
 3. How many pounds of chlorine will be required to make a ton 
 of bleaching powder? 
 
 4. How much manganese dioxide will be required to generate the 
 chlorine? 
 
 5. How much hydrochloric acid containing 40 per cent of the 
 acid will be required to generate the chlorine? 
 
 6. How much salt and how much sulfuric acid will be required to 
 give the hydrochloric acid? 
 
CHAPTER IX 
 
 GROUP VII: THE HALOGEN FAMILY, BROMINE, 
 IODINE AND FLUORINE 
 
 Groups of Elements. Although the vast number of 
 compounds which are known, with their infinite variety of 
 compositions and properties, are all made by different 
 combinations of eighty or a few more elements, even these 
 elements are not substances which are wholly different from 
 each other. There are three elements, bromine, iodine and 
 fluorine, which closely resemble chlorine. In a similar 
 manner there are three elements with properties somewhat 
 resembling those of oxygen, and the other elements fall into 
 a series of well-defined groups. After studying carefully 
 the properties of one of the elements of such a group it is 
 found that many of these properties are repeated with com- 
 paratively few differences in the other members of the 
 same group. 
 
 . Halogen Family. The group of elements which includes 
 chlorine is called the halogen family. The word halogen 
 means salt former and the name was given because the ele- 
 ments combine directly with metals to form salts which 
 contain only two elements, metal and halogen. Common 
 salt, or sodium chloride, is an illustration. The name con- 
 trasts these salts with the more common salts, such as 
 sodium sulfate, Na 2 SC>4, which contain oxygen in addition 
 to the metal and the non-metal. 
 
 The elements of the halogen family are : 
 
 Fluorine, 19 
 Chlorine, 35.5 
 Bromine, 80 
 Iodine, 127 
 
 92 
 
BROMINE 93 
 
 The atomic weights are given in round numbers. The 
 student is advised to learn these atomic weights, as the 
 atomic weights of the succeeding groups of non-metallic 
 elements are closely related to these and these numbers are 
 the best method which has been found for grouping the 
 elements and retaining a knowledge of the relations between 
 them in mind. 
 
 Bromine. Sea water and most of the brines from which 
 salt, NaCl, is obtained by evaporation contain small 
 quantities of sodium bromide, NaBr, and still smaller 
 quantities of sodium iodide, Nal. The electrolysis of such 
 a brine by passing an electric current through it causes 
 the liberation of iodine first, then of bromine and finally 
 chlorine. The amount of iodine in most of the brines is 
 very small indeed, but some of the brines, in Michigan, 
 especially, contain enough bromine so that the element may 
 be extracted profitably from them. After passing an 
 electric current through the brine till all of the bromine and 
 a little chlorine are liberated the bromine is distilled away. 
 Bromine is a dark liquid about three times as heavy as 
 water. It gives off a heavy, reddish brown vapor which is 
 very irritating and poisonous and which has a disagreeable 
 odor. In fact the name bromine is derived from a Greek 
 word, jftpwjuos, meaning a stench. 
 
 Compounds of Bromine. Bromine forms many com- 
 pounds which are similar to the corresponding compounds 
 of chlorine. Among these may be mentioned hydrobromic 
 acid, HBr, sodium bromide, NaBr, potassium bromide, 
 KBr, hypobromous acid, HBrO, potassium hypobromite, 
 KBrO, bromic acid, HBrO 3 , and potassium bromate, 
 KBrO,. 
 
 Frofti the analogy with chlorine answer the following 
 questions: 
 
 What will be the effect of water on hydrobromic acid 
 gas? 
 
4 THE HALOGEN FAMILY 
 
 Write the equations for the reaction between hydro- 
 bromic acid and manganese dioxide. Also for that between 
 hydrobromic acid and potassium permanganate. 
 
 Write the equations for the reaction by which potassium 
 hypobromite is prepared. Also the one for the reaction by 
 Which potassium bromate is formed. 
 
 What will be formed when potassium bromate is 
 heated? 
 
 Sodium Bromide, NaBr, is used in medicine to induce 
 sleep. Potassium bromide, KBr, which was used formerly 
 for the same purpose, is less suitable because of the irritating 
 effect of the potassium which it contains. 
 
 Silver Bromide, AgBr, is used in photography (see p. 311). 
 
 Iodine.-^-The occurrence of small amounts of iodine in 
 sea water and in brine? has been mentioned. It is not 
 practicable to obtain the element directly from these 
 Sources, but some sea weeds accumulate iodine in their 
 growth and from the ash of these weeds, called kelp, the 
 iodine can be obtained. 
 
 Free iodine crystallizes in black, shining scales. It 
 hielts when heated gently and gives off a beautiful violet 
 Vapor. The vapors condense as crystals on any cool 
 Surface, An evaporation of a solid without melting and 
 condensation of the vapor is called sublimation. Iodine 
 melts below its boiling point, but it also sublimes below 
 its melting point. 
 
 Iodine dissolves only very slightly in water. The solution 
 in alcohol is called tincture of iodine. It is used as an 
 application to bruises and swellings because of its germicidal 
 properties. The solution of iodine in alcohol is brown, but 
 that in carbon disulfide and in some other solvents is violet. 
 With a solution of starch in the presence of a little hydriodic 
 acid or an iodide, iodine gives an intense blue color. This 
 may be used as a very sensitive test either for iodine or 
 starch. 
 
IODINE. FLUORINE 95 
 
 Iodine forms many compounds similar to those of bromine 
 and chlorine. Name and give the formulas of some of 
 those which might be expected. 
 
 Potassium Iodide, KI, is used in medicine. An organic 
 compound containing iodine is found in the thyroid gland 
 of sheep and the lack of this compound in the human body 
 seems to be a chief cause of goiter. This compound has 
 been prepared, recently, at the Mayo laboratory, Rochester, 
 Minnesota, and gives remarkable results when used as a 
 medicine. 
 
 Fluorine. The compounds of chlorine, bromine and 
 iodine with calcium, calcium chloride, CaCl 2 , calcium 
 bromide, CaBr 2 , and calcium iodide, Cal2, are all easily 
 soluble in water, but the corresponding compound of fluor- 
 ine, CaF 2 , is almost insoluble. As compounds of calcium 
 are present in practically all natural waters it is impossible 
 for sea water or for brines to contain more than a very 
 small amount of fluorine. The element is found chiefly 
 in the form of calcium fluoride, CaF 2 , as the mineral fluorite 
 or fluorspar. The mineral crystallizes in cubes (or octa- 
 hedra). The pure mineral is colorless but it is often colored 
 violet or amethyst by the presence of a very minute amount 
 of some other compound, or substance, possibly a com- 
 pound of manganese. 
 
 Properties of Fluorine. Fluorine is the most active of 
 the non-metallic elements. While chlorine and other 
 elements of the family can be easily prepared by the 
 electrolysis of solutions of their compounds or by easy 
 reactions which occur in the presence of water, the electro- 
 lysis of an aqueous solution of a fluoride gives oxygen or 
 ozone in place of fluorine because free fluorine reacts with 
 water, liberating ozone, Os : 
 
 3F 2 + 3H 2 O = 6HF + O 3 
 
 Hydro- Ozone 
 fluorio 
 acid 
 
96 THE HALOGEN FAMILY 
 
 It was not till this property of fluorine was clearly 
 recognized that the French chemist, Moissan, was able to 
 prepare the free element by the electrolysis of anhydrous 
 hydrofluoric acid, HF, at a low temperature. He added a 
 small amount of potassium fluoride, KF, to the acid because 
 the pure acid is almost a non-conductor of electricity. By 
 electrolyzing the solution of potassium fluoride in anhydrous 
 hydrofluoric acid in a U-tube of platinum or of copper at a 
 temperature of 80 or below, hydrogen is liberated at the 
 cathode and fluorine at the anode. 
 
 At ordinary temperatures fluorine is a gas of a light, 
 greenish yellow color, similar to the color of diluted chlorine. 
 It combines directly with almost all other elements except 
 oxygen. Amorphous carbon, silicon, phosphorus and some 
 other elements take fire and burn when brought in contact 
 with fluorine. In each case a fluoride is formed, carbon 
 tetrafluoride, CF 4 , silicon tetrafluoride, SiF 4 , and phos- 
 phorus pentafluoride, PF 5 . 
 
 With the exception of the elements of the argon family, 
 which do not combine with other elements at all, fluorine 
 is the only element which does not combine with oxygen. 
 
 Compounds of Fluorine. Etching Glass. Hydrofluoric 
 acid, HF or H 2 F 2 , is prepared by the action of concentrated 
 sulfuric acid on calcium fluoride on the same principle as 
 the preparation of hydrochloric acid from salt: 
 
 NaCl + H 2 SO 4 =>NaHSO 4 + HC1 
 CaF 2 + H 2 SO 4 <=CaSOi + 2HF 
 
 In each case the acid, hydrochloric acid or hydrofluoric 
 acid, escapes from the mixture because it is a gas and the 
 disturbance of the equilibrium causes the reaction to go 
 always in the direction to form more of the compound 
 which escapes. 
 
 One of the most interesting and important properties 
 of hydrofluoric acid is its effect on glass. Ordinary glass 
 
COMPARISON 97 
 
 consists mainly of a mixture of silicates of calcium and so- 
 dium. When hydrofluoric acid gas comes in contact with 
 it the fluorine unites both with the silicon and with the 
 metals, forming silicon tetrafluoride, SiF 4 , calcium fluoride, 
 * CaF 2 , and sodium fluoride, NaF. The hydrogen of the 
 acid combines with the oxygen of the silicate to form water. 
 The silicon tetrafluoride, SiF 4 , is a gas and escapes. Hydro- 
 fluoric acid does not affect wax or paraffin. If a piece of 
 glass is covered with paraffin and a portion of the surface 
 of the glass is exposed by drawing lines or figures through 
 the wax, on exposing the object to the action of hydrofluoric 
 acid gas the parts of the glass uncovered will be etched. 
 In this manner the scales of thermometers and of other 
 instruments are made. In some cases solutions or pasty 
 mixtures containing hydrofluoric acid are used in place of the 
 gas. 
 
 Comparison of the Elements of the Halogen Family. 
 The colors of the halogens grow deeper with increasing 
 atomic weight. Fluorine is light greenish yellow; chlorine 
 is similar, but of a deeper shade; bromine vapor is reddish 
 brown, and iodine vapor is violet. Solid iodine is black. 
 
 At room temperature fluorine and chlorine are gases; 
 bromine is a liquid, iodine is a solid. 
 
 All of the family combine with hydrogen and the metals. 
 The following compounds may be taken as illustrations: 
 
 HF HC1 HBr HI 
 
 NaF NaCl NaBr Nal 
 
 CaF 2 CaClo CaBr 2 CaI 2 
 
 A1F 3 A1C1 3 AlBr 3 A1I 3 
 
 In these compounds with hydrogen and the metals the 
 affinity of the halogens decreases with increasing atomic 
 weight. Hydriodic acid, HI, undergoes considerable de- 
 composition if heated to 300 or 400, hydrobromic acid 
 is much more stable at that temperature, while hydro- 
 
98 THE HALOGEN FAMILY 
 
 chloric and hydrofluoric acids undergo scarcely any de- 
 composition, even at 1000. Any element of the family 
 will take hydrogen or a metal away from any other member 
 of the family having a higher atomic weight. Fluorine 
 will liberate chlorine from hydrochloric acid, bromine from 
 hydrobromic acid, or iodine from hydriodic acid. Bromine 
 will liberate iodine from hydriodic acid or from potassium 
 iodide, KI. 
 
 In the compounds with oxygen the affinity of the elements 
 increases with increasing atomic weight. Fluorine does 
 not combine with oxygen; chlorine monoxide, CljO, and 
 chlorine peroxide, C1O 2 , explode when heated, while iodine 
 pentoxide, I2O 5 , is a comparatively stable solid. 
 
 The oxygen acids increase in stability with an increase in 
 the amount of oxygen. Perchloric acid, HC104, is much 
 more stable than chloric acid, HC10 3 , and chloric acid is 
 more stable than hypochlorous acid. 
 
 SUMMARY 
 
 There are several groups of closely related elements. 
 
 The halogen family (Group VII) consists of fluorine, 
 chlorine, bromine and iodine. 
 
 Bromides are found in brines, with chlorides. 
 
 Bromine forms bromides, hypobromites and bromates. 
 
 Sodium bromide is much used in medicine ; silver bromide 
 in photography. 
 
 Iodides are found in sea water and in the ash of sea weeds. 
 
 Iodine forms iodides, iodates and periodates. 
 
 Iodine dissolves in alcohol. It gives a dark blue color 
 with starch. 
 
 Tincture of iodine, potassium iodide and an organic 
 compound of iodine from the thyroid gland are used in 
 medicine. 
 
 Fluorine occurs in calcium fluoride, or fluorspar. 
 
EXERCISES. HALOGENS 99 
 
 Fluorine is prepared by the electrolysis of anhydrous 
 hydrofluoric acid containing potassium fluoride. It is 
 the most active of the non-metallic elements. 
 
 Hydrofluoric acid is prepared by the action of sulfuric 
 acid on fluorspar. It is used in etching glass. 
 
 The halogens have a negative valence (pp. 104 and 106) 
 of one when combined with hydrogen or the metals. In 
 these compounds each will displace any of those of higher 
 atomic weight. 
 
 Fluorine does not combine with oxygen. The stability of 
 the compounds of the other halogens with oxygen increases 
 with increasing atomic weight and with increasing amounts 
 of oxygen. 
 
 EXERCISES 
 
 1. One liter of hydrobromic acid gas weighs approximately 
 3.62 grams. How many grams of sodium bromide would be 
 required to prepare 22.4 liters of the gas if it could be obtained 
 by the same method which is used for the preparation of hydro- 
 chloric acid? 
 
 2. How many grams of sulfuric acid will be required for the 
 same preparation? 
 
 3. Sodium bromide, manganese dioxide and sulfuric acid 
 interact with each other, giving bromine, manganese sulfate, 
 MnS0 4 , sodium sulfate, Na 2 S0 4 , and water. How many grams 
 of each compound will be required to give 160 grams of bromine? 
 
 4. A liter of hydrogen weighs approximately 0.09 gram. How 
 many liters of hydrogen will be given by the action of an excess of 
 hydrochloric acid in the following: 39 grams of potassium, 56 
 grams of iron, 27 grams of aluminium, 1 18 grams of tin (giving 
 SnCl 2 ), 65 grams of zinc? 
 
CHAPTER X 
 CLASSIFICATION OF ELEMENTS. VALENCE 
 
 Other Families of Elements. The elements of Group VI, 
 the sulfur family, with one exception, have atomic weights a 
 little less than those of the chlorine family. The elements of 
 the nitrogen family, in turn, have atomic weights a little 
 less than those of the sulfur family, and the atomic weights of 
 the carbon family are less than those of the nitrogen family. 
 The elements of the family of " Noble gases," which have no 
 chemical affinity, have atomic weights a little greater than 
 those of the chlorine family. The following table gives 
 boron in addition to the families, or groups of elements 
 mentioned. It includes all of the non-metallic elements and 
 a few half-metals and metals. This table furnishes a very 
 convenient basis for remembering the properties and rela- 
 tions of these elements and the student is advised to memo- 
 rize the atomic weights of the elements of each family as 
 the successive families are studied. 
 
 Group III 
 
 Group IV 
 
 Group V 
 
 Group VI 
 
 Group VII 
 
 Group O 
 
 
 
 
 
 
 He 4 
 
 Bll 
 
 C12 
 
 N14 
 
 O 16 
 
 F19 
 
 Ne20 
 
 A127 
 
 Si 28 
 
 P31 
 
 S32 
 
 Cl 35.5 
 
 A 40 
 
 Ga 69.9 
 
 Ge 72.5 
 
 As 75 
 
 Se79 
 
 Br80 
 
 Kr83 
 
 In 114.8 
 
 Sn 118 
 
 Sbl20 
 
 Te 127.6 
 
 1127 
 
 Xel30 
 
 T1204 
 
 Pb 207 
 
 Bi 208 
 
 
 
 Nt222 
 
 
 
 
 
 
 
 A study of these elements has made it very clear that 
 there is a close .connection between the properties of an 
 
 100 
 
VALENCE A 
 
 element and its atomic weight. The table is part of a 
 larger table in which all of the elements are arranged in 
 accordance with their atomic weights. This larger table, 
 which will be considered later (p. 162), is used as the basis 
 for the classification of the elements. 
 
 Valence. The atomic weight of sodium is 23, that of 
 magnesium is 24 and that of aluminium is 27. 1 In accord- 
 ance with the atomic theory, if we take 23 milligrams of 
 sodium, 24 milligrams of magnesium and 27 milligrams of 
 aluminium, we shall have the same number of atoms of each 
 
 FIG. 27. 
 
 element. If we allow these quantities of the elements to 
 act on hydrochloric acid, we shall get from the sodium about 
 11 cc. of hydrogen gas, from the magnesium 22 cc. and from 
 the aluminium 33 cc. (Fig. 27) . It is evident from this that 
 one atom of magnesium displaces twice as much hydrogen 
 and combines with twice as much chlorine as one atom of 
 sodium; and that one atom of aluminium displaces and 
 combines with three times as much hydrogen and chlorine 
 
 1 These atomic weights are selected on the basis of the specific heats of 
 the elements (see p. 266.) 
 
102 CLASSIFI CATION OF ELEMENTS 
 
 as the sodium atom. These relations are clearly expressed 
 in the following equations: 
 
 Na + HC1 = NaCl + H 
 Mg + 2HC1 = MgCl 2 + 2H 
 Al + 3HC1 = A1C1, + 3H 
 
 Similar differences between the non-metallic elements are 
 illustrated by the following tables of their compounds with 
 hydrogen : 
 
 CH 4 NH 3 OH 2 FH 
 
 SiH 4 PH 3 SH 2 C1H 
 
 AsH 3 SeH 2 BrH 
 
 SbH 3 TeH 2 IH 
 
 A study of these and other compounds has led to the 
 conclusion that in chemical compounds each atom unites 
 directly with only a small number of other atoms and that 
 this number of atoms with which an atom combines varies 
 for the different elements. This property is called valence. 
 Atoms, such as those of sodium or chlorine, which combine 
 with a single atom of chlorine or hydrogen, are said to have a 
 valence of one and are called univalent. Those which unite 
 with two atoms, such as the atoms of magnesium and of 
 oxygen, have a valence of two and are called bivalent. 
 Aluminium and phosphorus have a valence of three and are 
 called trivalent. Carbon and silicon have a valence of 
 four and are called quadrivalent. Oxygen is evidently 
 bivalent in water, H 2 O, and it seems to be bivalent in nearly 
 or quite all of its stable compounds. Taking this into con- 
 sideration in the series 
 
 Na 2 0, MgO, A1 2 O 3 , SiO 2 , P 2 5 , S0 3 , C1 2 O 7 
 
 the valence of the elements increases from one to seven and 
 the elements are called, successively, univalent, bivalent, 
 trivalent, quadrivalent, quinquivalent, sexivalent and 
 septivalent. 
 
GRAPHICAL FORMULAS 103 
 
 Varying Valences. The valence of elements is not, 
 however, so simple as the statements of the previous para- 
 graph would seem to imply. Phosphorus has a valence of 
 three in the compound PH 3 , but it forms two compounds, 
 PC1 3 and PC1 5 , with chlorine. In the first of these it is 
 trivalent but in the second it is quinquivalent. 
 
 In such a series of compounds as the following: 
 
 N 2 0, NO, N 2 3 , N0 2 , N 2 5 
 nitrogen seems to vary in its valence from one to five. 
 
 While it is clear from the illustrations given that the same 
 element may have a different valence in different compounds 
 the fundamental idea of valence, that in compounds each 
 atom holds directly to only a small, definite number of 
 other atoms, has proved very useful. 
 
 Graphical Formulas. The valence of elements in com- 
 pounds may be clearly expressed by formulas in which the 
 symbols of the elements are taken to represent atoms, and 
 lines are used to represent the direct attachments which 
 exist between the atoms. This gives such formulas as the 
 following, which will be easily understood without further 
 explanation. 
 
 /Cl /Gl 
 
 Na - Cl MgQ Al^Cl 
 
 X C1 X C1 
 
 H H 
 
 I / 
 
 H - C - H N H H-O-H H-Cl 
 
 I \ 
 
 H H 
 
 J3 ,,0 
 
 M g =o \> off No s^o ;o 
 
104 CLASSIFICATION OF ELEMENTS 
 
 Such formulas are called graphical formulas. 
 
 Groups of elements may also have a definite valence, as 
 is seen in such compounds as sodium sulfate, Na2S0 4 , and 
 sodium phosphate, Na 3 PO 4 , in which the group SO 4 is 
 bivalent and the group P0 4 trivalent. This is more clearly 
 expressed by the formulas: 
 
 Na, 
 
 N 
 N 
 Na' Na 
 
 In all such cases the valences of the two parts must agree 
 and this gives valuable assistance in the writing of correct 
 formulas. 
 
 It is supposed, of course, that the valence of the group 
 is dependent, ultimately, on the valences of the atoms of 
 which it is composed. This may be expressed by such 
 formulas as: 
 
 Na - Ov ,> Na - (\ 
 
 )S^ and Na - (A? = O 
 
 Na - O/ ^O Na - O / 
 
 Positive and Negative Valences. In the electrolysis 
 of hydrochloric acid, HC1, sodium chloride, NaCl, or of other 
 metallic chlorides, the chlorine goes to the anode or positive 
 pole while the hydrogen or metal goes to the cathode or 
 negative pole. The same is true of hydrogen and metallic 
 compounds of the other halogens, fluorine, bromine and 
 iodine. For this reason hydrogen and the metals are 
 called positive because they are attracted by the negative 
 electrode, and the halogens are called negative because 
 they are attracted by the positive electrode. In the 
 electrolysis of dilute sulfuric acid, or of a solution of 
 sodium hydroxide, hydrogen is liberated at the cathode 
 and oxygen at the anode and this gives us a good reason for 
 
POSITIVE AND NEGATIVE VALENCES 105 
 
 considering oxygen negative. On account of these relations 
 we may say that hydrogen and sodium have a positive 
 valence of one; calcium has a positive valence of two; 
 chlorine and bromine have a negative valence of one, and 
 oxygen has a negative valence of two. 
 
 If some elements were always positive and others negative 
 it would be an easy matter to classify elements on this basis. 
 But some elements, such as nitrogen, sulfur and many 
 others, combine with both hydrogen and with oxygen. 
 We have, for instance, 
 
 < 
 
 N^H and )O 
 
 H < 
 
 
 
 Ammonia Nitrogen trioxide 
 
 We are compelled either to believe that nitrogen has a 
 negative valence in ammonia and a positive valence in 
 nitrogen trioxide, or that the positive or negative character 
 of an element disappears when elements unite. It is only 
 fair to say that chemists are not altogether certain which 
 of these alternatives is true. It is a question under debate 
 and the relations are so complicated that a satisfactory 
 decision may not be reached for some time. Meanwhile 
 the distinction between positive and negative valences is 
 one of increasing importance in the study of chemistry. 
 In the following table positive valences are indicated by the 
 plus sign and negative valences by the minus sign. 
 
 A study of the table shows that positive valences are much 
 more common than negative, and this is true even of the 
 non-metallic elements. No element with an atomic weight 
 above 130 shows a negative valence. It is also doubtful 
 if any element of the first three groups, except boron, 
 shows a negative valence. 
 
106 
 
 CLASSIFICATION OF ELEMENTS 
 
 
 oo -a 
 
 ~t . 
 O co" _c 
 
 $ 
 
 I ^ 
 
 + 
 
 M 
 
 | 
 
 + + + 
 
 1 
 
 iO ^ 
 
 i+ 
 
 IO 
 
 ?SJ: 
 
 ? 
 
 
 
 + 
 
 
 
 o+ 
 
 CO C3 CO ^H + 
 + 
 
 g,<N 
 + 
 
 C (M ^,<M -O 
 
 tSJ + 02+ 
 
 M "T 
 
 o 
 
 5? 
 
 
 
 a 
 
SUMMARY. VALENCE 107 
 
 The following table gives the most common valences for 
 each group: 
 
 Common Valences 
 
 Group 
 
 Group I + 1 
 
 Group II + 2 
 
 Group III -f 3 
 
 Group IV +4, - 4 
 
 Group V + 3, -h 5, - 3(- 4 + 1) 
 
 Group VI + 2, + 4, + 6, - 2(- 3 + 1) 
 
 Group VII + 1, + 3, + 5, + 7,-l(- 2 + 1) 
 
 Group VIII + 2, + 3', + 4, + 8 
 
 SUMMARY 
 
 A very satisfactory classification of the five families of 
 non-metallic elements is based on their atomic weights. , 
 
 An atom of each element, except those of the zero group, 
 has the power to combine directly with a definite, small 
 number of other atoms. This power is called valence. 
 Valences vary both for different elements and for the same 
 element in different compounds. 
 
 The distribution of the valences in compounds may be 
 most clearly expressed by graphical formulas. 
 
 Some valences, as those of hydrogen and the metals, 
 are positive; others, as those of the halogens in the halides, 
 are negative. 
 
 Elements with positive valences are much more common 
 than those with negative valences. 
 
 A table is given for the valences of the elements. 
 
 EXERCISES 
 
 Write graphical formulas for the following compounds : manga- 
 nese dioxide, manganese chloride, manganese sulfate, potassium 
 hypochlorite, potassium chlorite, potassium chlorate, potassium 
 perchlorate, potassium permanganate. In the salts represent 
 
108 CLASSIFICATION OF ELEMENTS 
 
 the potassium as united to the other elements through one oxygen 
 atom. 
 
 While the graphical formulas of inorganic compounds which are 
 referred to here have considerable value and very probably re- 
 present the real structure of these compounds, too much value 
 should not be attached to them. Especially, it is important to 
 understand that valence alone is not a sufficient guide for the writing 
 of graphical or structural formulas. The two formulas 
 
 ^ 
 H- 0-0-0-0- Cl and H - - Cl =0 are both in ac- 
 
 ^0 
 
 cord with the valence of chlorine in other compounds but only 
 one of these can be true and we are not altogether certain that 
 either is true, though there is considerable probability of the truth 
 of the second. 
 
CHAPTER XI 
 GROUP VI: SULFUR, SELENIUM AND TELLURIUM 
 
 Occurrence. Chlorine and the other elements of the 
 chlorine family are not found free in nature, chiefly because 
 of their strong affinity for metals which are common and 
 abundant. Oxygen, on the other hand, is found free partly 
 because of the great abundance of the element and partly 
 because the growth of plants constantly liberates oxygen 
 from the carbon dioxide of the air. Sulfur, the second 
 element of the oxygen group, is also found in the free state 
 largely because oxygen liberates sulfur from hydrogen 
 sulfide very much as fluorine would liberate chlorine from 
 hydrochloric acid or as chlorine liberates bromine from 
 bromides. 
 
 Free sulfur is found especially in Sicily and in Louisiana. 
 For a long time the sulfur mines of Sicily held a practical 
 monopoly of the sulfur markets of the world. The sulfur 
 of Louisiana is beneath a bed of quicksand and for many 
 years after the deposits were discovered no one was able 
 to devise a method for mining the mineral. Finally 
 Hermann Frasch, of New York, conceived the idea that it 
 would be possible to melt the sulfur by means of hot water 
 under pressure and he devised a system of concentric tubes 
 by means of which hot water is brought down on top of the 
 sulfur and the melted sulfur is brought to the surface by 
 means of a current of air. This process now furnishes all 
 of the free sulfur required for manufacturing purposes in 
 the United States. 
 
 In Sicily the sulfur is mixed with other minerals from 
 
 109 
 
110 SULFUR, SELENIUM AND TELLURIUM 
 
 which it is separated by melting it. The heat is furnished 
 sometimes by burning a part of the sulfur, as other fuels 
 are scarce and expensive. The sulfur is further purified by 
 distilling it. If the vapors are condensed in warm chambers 
 so that the sulfur remains liquid, it is cast in rolls and called 
 roll brimstone. If the condensing chambers are cold so 
 that the vapor condenses at once to a solid the product is 
 called flowers of sulfur. The difference between the for- 
 mation of roll brimstone and flowers of sulfur is similar to 
 that between the condensation of water vapor from the 
 clouds as rain and as snow. 
 
 Sulfur is also found in nature in the form of sulfides and 
 of sulfates. Several important ores of common metals are 
 sulfides. Among these may be mentioned galena, or lead 
 sulfide, PbS, sphalerite, or zinc sulfide, ZnS, and chalcopy- 
 rite, or copper pyrites, CuFeS2. Iron pyrites or pyrite, 
 FeS 2 , is extensively used as a source of sulfur for the manu- 
 facture of sulfuric acid, and is to be considered as an ore 
 of sulfur rather than as an ore of iron. The most important 
 sulfates are calcium sulfate, found as the hydrate, CaSO 4 . 
 2H 2 O, called gypsum or alabaster and used in the manu- 
 facture of plaster of Paris, and barium sulfate, BaSO 4 , used 
 in making compounds of barium and as a substitute for 
 white lead. 
 
 Sulfur is used extensively in "sulfuring fruit" which is to 
 be dried, the fruit being exposed to the action of sulfur 
 dioxide obtained by burning the sulfur. The sulfur diox- 
 ide kills microorganisms which would injure the fruit before 
 it is dry and also prevents the darkening of the fruit. 
 
 Sulfur is also dissolved in milk of lime to give a "lime- 
 sulfur wash" which is applied to vines and fruit trees to 
 prevent the growth of fungi and other harmful organisms. 
 
 Sulfur is a constituent of the old-fashioned black gun- 
 powder but is not used in the smokeless powders. It is used 
 in the manufacture of india-rubber and of carbon disul- 
 
FORMS OF SULFUR 111 
 
 fide and in some dyes, especially in sulfur black. It is 
 sometimes used in making sulfuric acid, but iron pyrites, 
 FeS 2 , is usually a cheaper source of sulfur for that purpose. 
 
 Forms of Sulfur. There are two allotropic forms of 
 oxygen, ordinary oxygen and ozone. Sulfur gives three 
 allotropic, solid forms. 
 
 The sulfur found in nature crystallizes in rhombic crystals 
 which melt at 114.5. Similar crystals can be prepared 
 by dissolving sulfur in carbon disulfide and allowing the 
 solution to evaporate. 
 
 If sulfur is heated till it melts and then is allowed to cool 
 slowly it crystallizes in long needles which can be obtained 
 by breaking the crust on the surface, when the mass is 
 partially solidified, and pouring out the liquid portion. 
 These needles are nearly transparent at first and melt 
 at 119. If they are allowed to cool and stand for some 
 time they become opaque and the melting point changes 
 to 114.5, showing that they have changed to microscopic 
 crystals of the rhombic variety. It is evident from this 
 that the rhombic crystals are the stable form at ordinary 
 temperatures while the needles are more stable at a tempera- 
 ture near the melting point. 
 
 When sulfur is heated above its melting point, at 160 
 it suddenly becomes thick and viscous. If it is cooled 
 slowly it changes back as it cools to the ordinary forms, 
 but if it is cooled quickly it may be obtained in a form which 
 is at first plastic like india-rubber but which may solidify 
 without crystallizing. This is called amorphous sulfur, 
 meaning that it has no crystalline form and it is to be con- 
 sidered as a supercooled liquid rather than as an ordinary 
 solid. This amorphous sulfur is insoluble in carbon 
 disulfide. 
 
 Liquid and Gaseous Sulfur. At a temperature slightly 
 above its melting point sulfur is a pale, yellow, mobile 
 liquid. At 160, as stated above, it suddenly becomes 
 
112 SULFUR, SELENIUM AND TELLURIUM 
 
 very viscous and the character of the substance obtained 
 by cooling this liquid shows that the sulfur changes to an 
 allotropic form at that temperature. When heated still 
 hotter the liquid becomes more mobile again, but remains 
 of the same dark color which is characteristic of the viscous 
 liquid. At 445 the sulfur boils. If care is taken to keep 
 the walls of the tube or flask above the boiling liquid at 
 a temperature above the boiling point it can be seen that 
 the vapor is of a light amber color. 
 
 Hydrogen Sulfide. Waters from artesian wells often 
 contain hydrogen sulfide in solution and such waters are 
 known 'as "sulfur" waters. The hydrogen sulfide gives 
 to the waters a disagreeable odor and taste. Hydrogen 
 sulfide is one of the substances which gives the offensive 
 odor of decayed eggs and it is usually present in sewer gas. 
 It is prepared by the action of hydrochloric or sulfuric 
 acid on ferrous sulfide, FeS, just as water is formed when 
 an acid acts on zinc oxide, ZnO: 
 
 ZnO + 2HC1 = ZnCl 2 + H 2 O 
 FeS + 2HC1 = FeCl 2 + H 2 S 
 
 Ferrous Ferrous 
 
 . sulfide chloride 
 
 Hydrogen sulfide burns in air to sulfur dioxide, SO 2 , 
 and water. If a test-tube or cylinder filled with the gas 
 is ignited at the mouth so that the gas burns with an in- 
 sufficient supply of air, the hydrogen burns to water while 
 a part of the sulfur is separated in the free state. 
 
 Sulfides. Groups of Metals for Analysis. Many of the 
 metals combine directly with sulfur when heated with it. 
 Many of the sulfides of metals are insoluble in water 
 and the addition of hydrogen sulfide to solutions of salts 
 of these metals will cause the sulfides to separate as pre- 
 cipitates. Thus if hydrogen sulfide is passed into an acid 
 solution containing lead, arsenic or antimony, a precipitate 
 
GROUPS OF QUALITATIVE ANALYSIS 113 
 
 of lead sulfide, PbS, arsenic sulfide, As 2 S 3 , or antimony 
 sulfide, Sb 2 S 3 , will separate. From a neutral or alkaline 
 solution containing zinc or iron, zinc sulfide, ZnS, or ferrous 
 sulfide, FeS, will separate. This conduct of solutions of the 
 metallic salts furnishes a ready means of separating the 
 metals into three groups for purposes of analysis. The 
 first group contains lead, arsenic, antimony and other 
 metals whose sulfides are so insoluble that they may be 
 precipitated from acid solutions by .hydrogen sulfide. 
 The second group consists of metals whose sulfides are 
 slightly more soluble and for that reason are not precipitated 
 from acid solutions but they are precipitated from alkaline 
 solutions. , 
 
 If an acid solution containing metals of all three groups 
 is treated with hydrogen sulfide, the sulfides of the first 
 group are precipitated and by filtering off and washing 
 the precipitates the metals of that group may be separated 
 from the metals of the second and third groups. By adding 
 ammonia to the filtrate, that is, to the solution containing 
 metals of the second and third groups which runs through 
 the filter, the second group will be precipitated .and this 
 precipitate may be filtered off and separated from the 
 metals of the third group. 1 
 
 Formation of Precipitates. In a previous chapter 
 (p. 65) the formation of a precipitate of silver chloride, 
 AgCl, has been explained as due to the presence of ions 
 which readily separate from each other or combine to form 
 new compounds in solution. An extension of the same 
 principle helps us to understand the precipitation of sul- 
 fides. For a better understanding of what occurs two other 
 principles are needed. These are: 
 
 First, that the formation of ions and the formation of 
 
 1 In the classification adopted in most books on qualitative analysis 
 the metals of Group I as given here are separated into two groups called 
 I and II; the metals of Group II above are given as Group III and Group 
 III above is divided iato Groups IV and V. 
 8 
 
114 SULFUR, SELENIUM AND TELLURIUM 
 
 precipitates and other compounds in solutions are reversible 
 reactions. 
 
 Second, that, as with all other reversible reactions, the 
 increase of one of the substances taking part tends to in- 
 crease the amount of the substances formed from it and a 
 decrease of the substance tends to decrease the amount 
 formed. 
 
 The precipitation of lead sulfide from a solution of lead 
 nitrate, Pb(NOs)2, depends on the reversible reactions rep- 
 resented by the following equations: 1 
 
 Pb(NO 3 ) 2 <=* Pb++ + N0 3 ~ + NOr 
 H 2 S <= S- + H+ + H+ 
 H+ + N0 3 - * HN0 3 
 Pb++ + S- * PbS 
 
 Combining these we have: 
 
 Pb(N0 3 ) 2 + H 2 S = PbS + 2HN0 3 
 4 
 
 Nitric acid ionizes to a much greater extent than hydrogen 
 sulfide, or, in other words, a solution of nitric acid contains 
 many hydrogen ions while one of hydrogen sulfide contains 
 relatively few of such ions. If we add nitric acid or any 
 other acid which gives a large number of hydrogen ions the 
 reversible reaction represented by the equation : 
 
 H 2 S <= S- + H+ + H+ 
 
 must be driven to the left by the increase in one of the con- 
 stituents, H + , of the right-hand side of the equation. This 
 will cause a decrease in the number of sulfur ions, S = , and 
 this may go so far as to cause the reaction represented by 
 the equation: 
 
 Pb++ + S- <= PbS 
 
 1 For the sake of simplicity the formation of intermediate ions such as 
 HS~ and Pb + SH and the hydrolysis of Na 2 S to NaHS and NaOH are 
 not considered here. See Noyes, Textbook of Chemistry, pp. 169, 170 
 
PRECIPITATION. SULFUR DIOXIDE 115 
 
 also to go to the left and the lead sulfide would dissolve. 
 It follows from this that while lead sulfide is precipitated 
 from a solution containing a dilute acid it will dissolve in a 
 more concentrated acid and the precipitate will not be 
 formed in the presence of much free acid. 
 
 The precipitation of ferrous sulfide from an alkaline solu- 
 tion depends on the fact that in the reaction represented by 
 the equation: 
 
 Na 2 S ?=* S- + Na+ + Na + 
 
 the equilibrium is much farther to the right and there are 
 many more sulfur ions, S = , present than in the correspond- 
 ing ionization of hydrogen sulfide. The larger number of 
 sulfur ions causes the formation of ferrous sulfide : 
 
 FeCl 2 <=* Fe++ + Cl~ + Cl~ 
 Fe++ + S- <F FeS 
 
 In all reactions when a precipitate is formed the compound 
 precipitated is removed from the sphere of action and this 
 drives the reaction in the direction toward its formation. 
 
 Hydrosulfides. When a hydroxide of a metal which is 
 not precipitated by hydrogen sulfide is treated with the gas 
 a hydrosulfide is formed : 
 
 NaOH + HSH = NaSH + HOH 
 
 The name hydrosulfide refers, of course, to a compound 
 containing hydrogen and sulfur, just as a hydroxide con- 
 tains hydrogen and oxygen. 
 
 Sulfur Dioxide. Sulfur burns in air or in oxygen with a 
 blue flame, giving a gas with a suffocating odor. This has 
 the composition represented by the formula S02 and is 
 called sulfur dioxide. The same odor is sometimes ob- 
 served from the cinders of coal containing pyrites. Sulfur 
 dioxide is a good germicide and its use in "sulfuring" fruit 
 has been spoken of. It was formerly used to disinfect 
 
116 SULFUR, SELENIUM AND TELLURIUM 
 
 rooms, clothing and articles which had been infected with 
 bacteria which cause infectious diseases, but its use for these 
 purposes has been almost entirely replaced by the use of 
 formaldehyde, which is even more effective and which does 
 not bleach fabrics and discolor metallic objects as sulfur 
 dioxide does. 
 
 Sulfur dioxide is used for bleaching, especially for bleach- 
 ing straw hats. It produces its effect by combining with the 
 coloring matter of the straw to produce a colorless com- 
 pound. These compounds are not very stable and they 
 are decomposed by the action of air and light, restoring 
 the original color of the straw. The action in this regard is 
 quite different from that of chlorine, which permanently 
 destroys the coloring matters which it attacks. Chlorine is 
 more liable to weaken the fiber of the material which is 
 treated, however. 
 
 Sulfur dioxide may be prepared by warming copper with 
 concentrated sulfuric acid. We may suppose that the 
 acid oxidizes the copper to cupric oxide, CuO, and is itself 
 reduced to sulfurous acid, H 2 SO 3 , and that the latter de- 
 composes into sulfur dioxide and water. The copper oxide 
 also reacts with more sulfuric aeid r giving copper sulfate r 
 CuS0 4 : 
 
 Cu + H 2 SO = [CuO]; -f H 2 S0 3 
 
 H 2 SO 3 = SO 2 -f H 2 O 
 .[CuO] + H 2 SO 4 = CuSO 4 -f H 2 O 
 
 Combining: 
 
 Cu + 2H 2 SO 4 = CuSO 4 + $O 2 + 2H 2 
 
 The copper oxide, CuO, is enclosed in brackets because 
 we have no positive evidence that it is formed as an inter- 
 mediate compound. We only know that the reaction takes 
 place in accordance with the last equation and that the 
 supposition that copper oxide is io-raied! temporarily is in 
 
SULFITES. DIBASIC ACIDS 117 
 
 accordance with the facts that sulfuric acid is reduced and 
 that the copper is oxidized in the reaction. 
 
 Sulfurous Acid. Sulfites. Water dissolves a certain 
 amount of sulfur dioxide. Part of the gas combines with 
 the water to form sulfurous acid, H 2 SO3, as is shown by the 
 acid reaction of the solution with litmus paper and by the 
 neutralization of bases by the solution. By passing sulfur 
 dioxide into a solution of sodium hydroxide, NaOH, or 
 milk of lime, Ca(OH) 2 , sodium sulfite, Na 2 SO 3 , or calcium 
 sulfite, CaS0 3 , may be prepared. The same salts may be 
 prepared by the action of sulfur dioxide and water on so- 
 dium carbonate, Na 2 CO 3 , or calcium carbonate, CaCO 3 
 because sulfurous acid dissociates (p. 150) less easily into 
 sulfur dioxide and water than carbonic acid dissociates to 
 carbon dioxide and water. 
 
 Acid Sulfites. Dibasic Acids. Sulfurous acid, H 2 SO 3 , 
 has two atoms of hydrogen in its molecules which may be 
 replaced by metals and it not only forms salts, such as sodium 
 sulfite, Na 2 SO 3 , and calcium sulfite, CaS0 3 , in which all of 
 the hydrogen Tias been replaced, but it also forms salts in 
 which only one-half of the hydrogen has been replaced. 
 Acid sodium sulfite, NaHS0 3 , and acid calcium sulfite, 
 CaH 2 (SO 3 ) 2 , are salts of this type. Such salts may be 
 prepared by adding to a solution of sulfur dioxide enough of 
 the base to neutralize only one-half of the acid: 
 
 H 2 S0 3 + NaOH = NaHSO 3 + H 2 O 
 
 The salts may also be prepared by passing sulfur dioxide 
 into a solution of the base till the odor of sulfur dioxide 
 becomes apparent. 
 
 2NaOH + SO 2 = Na 2 SO 3 + H 2 O 
 Na 2 SO 3 + SO 2 + H 2 O = 2NaHSO 3 
 
 Note. Similar reactions should be written for the prepa- 
 ration of the calcium salt. 
 
118 SULFUR, SELENIUM AND TELLURIUM 
 
 Acid sodium sulfite is a germicide and it has sometimes 
 been added to cider or wine to stop the fermentation. The 
 acid calcium salt is used in paper making to dissolve un- 
 desirable substances from the fiber of the wood which is 
 used to make paper pulp. 
 
 The most convenient laboratory method for the prepara- 
 tion of sulfur dioxide is to drop concentrated sulfuric acid 
 into a 40 per cent solution of acid sodium sulfite. Write 
 the equation for the reaction. 
 
 Acids like sulfurous acid, which form two classes of salts 
 (here NaHSO 3 and Na 2 SO 3 ), in one of which only half of the 
 hydrogen is replaced, are called dibasic. In contrast with 
 these, hydrochloric acid, HC1, and nitric acid, HNO 3 , 
 are monobasic. Phosphoric acid, H 3 PO 4 , which forms the 
 salts Na 3 P0 4 , Na 2 HP0 4 and NaH 2 PO 4 , is tribasic. A salt 
 in which a metal has replaced all of the replaceable hydro- 
 gen of an acid is called a normal salt. Normal salts are not 
 always neutral (see p. 77). 
 
 Sulfur Trioxide or Sulfuric Anhydride. Catalysis. 
 Sulfur dioxide will combine directly with more oxygen to 
 form sulfur trioxide, S0 3 , but the reaction is a reversible 
 one: 
 
 2SO 2 + O 2 * 2SO 3 
 
 At ordinary temperatures the equilibrium in this reac- 
 tion is far toward the right, that is, when the equilibrium is 
 reached a large part of the sulfur dioxide and oxygen will 
 have combined to form sulfur trioxide and only a small 
 part will remain uncombined. But the combination takes 
 place so slowly that the equilibrium is reached only after a 
 very long time and the reaction could not be made commer- 
 cially profitable. 
 
 At a high temperature this combination takes place 
 rapidly but the equilibrium shifts more and more to the left, 
 and when equilibrium is reached a large part of the sulfur 
 
SULFUR TRIOXIDE HO 
 
 dioxide and oxygen remain uncombined. It is impossible 
 to find any temperature at which the combination is both 
 rapid enough and complete enough to be used as a commer- 
 cial process. 
 
 The practical solution of a difficulty of this kind is to find, 
 if possible, some catalytic agent (p. 9) which will increase 
 the speed of the reaction so that it may be carried on rapidly 
 at a temperature low enough so that the equilibrium is still 
 favorable to the production of the compound desired . It has 
 been known for more than seventy years that finely divided 
 platinum is a suitable catalytic agent to promote the union 
 of sulfur dioxide and oxygen, but many difficulties were met 
 in the practical details of the process and it was not till 
 about 1900 that these were so far overcome as to make the 
 "contact process" for the manufacture of sulfur trioxide 
 and sulfuric acid successful on a large scale. At the present 
 time large quantities of these compounds are manufactured 
 by this process. 
 
 Properties of Sulfur Trioxide. Pure sulfur trioxide is a 
 volatile liquid which boils at 46. A trace of moisture 
 causes it to unite with itself to form a solid compound, 
 S 2 O 6 , which is called a polymeric form. A polymer is a 
 compound formed by the union of two or more molecules 
 of the same kind. It has exactly the same proportions of 
 the elements as the original compound. We may suppose 
 that sulfur trioxide and its polymer have some such struc- 
 tures as the following: 
 
 ,0 % /Qv .0 
 
 0-< and X )< 
 
 ^O O^ X CK X) 
 
 The preparation of sulfur trioxide on a small scale may 
 be carried out with the apparatus shown in Fig. 28. A 
 mixture of sulfur dioxide and oxygen, dried by passing the 
 gases through bottles containing concentrated sulfuric 
 
120 
 
 SULFUR, SELENIUM AND TELLURIUM 
 
 acid, is passed through a tube containing platinized asbestos. 
 The tube is warmed gently and at a temperature of 300 
 to 400 the combination takes place rapidly. 
 
 Sulfur trioxide hisses like hot iron when thrown into water. 
 It combines with the water to form sulfuric acid. Because 
 of this relation to sulfuric acid it is often called sulfuric 
 anhydride, meaning sulfuric acid without water or from 
 which the water has been removed. 
 
 FIG. 28. 
 
 Sulfur trioxide fumes strongly in the air because it com- 
 bines with the moisture of the air to form sulfuric acid. 
 
 Effects of Temperature upon Reversible Reactions. 
 Principle of van't Hoff-Le Chatelier. It has been pointed 
 out that an increase of temperature hastens both the com- 
 bination of sulfur dioxide and oxygen and the decomposi- 
 tion of sulfur trioxide, but that the rate of decomposition 
 increases more rapidly than the rate of combination. In 
 general it has been found that the speed of any reaction is 
 doubled by an increase of 10 in the temperature. For an 
 increase of 20 the speed is four times as great, for 30 eight 
 times as great, and so on. But the fact that the two direc- 
 tions of a reversible reaction are not hastened at the same 
 rate shows that the statement just given is only a rough, 
 experimental rule. 
 
 Heat is generated when sulfur dioxide and oxygen unite. 
 It follows that heat is absorbed when sulfur trioxide is 
 decomposed. The effect of temperature on the equilibrium 
 

 LEAD CHAMBER PROCESS 121 
 
 is intimately connected with these facts. An increase in 
 temperature always opposes that direction of a reversible 
 reaction which causes an evolution of heat. To take another 
 illustration: oxygen and hydrogen combine with an evolu- 
 tion of heat. At high temperatures an increase in tempera- 
 ture causes more of the oxygen and hydrogen to remain 
 uncombined or, what is the same thing, causes more of the 
 water to decompose. 
 
 An increase in temperature opposes the combination of 
 the elements because their combination generates heat and 
 the heat generated tends to decompose the water formed. 
 
 The effect of an increase in temperature upon the com- 
 bination of sulfur dioxide and oxygen is an illustration of 
 the principle of van't Hoff-Le Chatelier, which is that 
 every force applied to a system in equilibrium produces a 
 change which tends to resist the force applied. If heat is 
 generated by the combination of two substances the 
 application of heat tends to decrease the amount of com- 
 bination because heat causes the decomposition of a part 
 of the product formed. 
 
 If we could understand fully the mechanism of all physical 
 changes it seems likely that we should find that the principle 
 has its foundation in Newton's law that for every action 
 there is an equal and opposite reaction. 
 
 Lead Chamber Process for Manufacturing Sulfuric 
 Acid. Many different oxidizing agents will oxidize sul- 
 furous acid, or sulfur dioxide and water, to sulfuric acid. 
 One of these, nitrogen dioxide, NO 2 , is reduced in the 
 process to nitric oxide, NO, a compound which combines 
 directly and rapidly with the oxygen of the air to form the 
 original compound, nitrogen dioxide. The reactions can 
 be expressed most simply by the following equations: 
 
 N0 2 + S0 2 + H 2 = H 2 S0 4 + NO 
 2NO + 2 =2NO 2 
 
122 SULFUR, SELENIUM AND TELLURIUM 
 
 In the light of these equations the nitrogen dioxide is 
 merely a catalytic agent which enables sulfur dioxide and 
 water to combine rapidly with the oxygen of the air to 
 form sulfuric acid. This method of making sulfuric acid 
 has been in use for one hundred and seventy years and is 
 known as the " chamber process." 
 
 The manufacture is carried out in large chambers built 
 of sheet lead. The sheets of lead are melted together at 
 the edges by means of the oxyhydrogen blowpipe, a process 
 known as "lead burning." and the chamber of lead is 
 supported by a. wooden framework on the outside. 
 
 The sulfur dioxide for the process is usually made by 
 burning iron pyrites, FeS 2 , in a furnace especially con- 
 structed for the purpose: 
 
 2FeS 2 + HO = 4SO 2 + Fe 2 O 3 
 or 
 
 4FeS 2 + 11O 2 = 8SO 2 + 2Fe 2 O 3 
 
 To furnish the nitrogen dioxide nitric acid, HN0 3 , is 
 introduced. This reacts at first with the sulfur dioxide 
 and water, giving sulfuric acid and nitric oxide: 
 
 2HNO 3 + 3SO 2 + 2H 2 = 3H 2 SO 4 + 2NO 1 
 
 Oxygen is furnished by the introduction of air, and water 
 is introduced in the form of a spray. 
 
 Since air is forced into the chambers the nitrogen of the 
 air must constantly escape at the end and this carries with 
 it the larger part of the nitrogen dioxide. To prevent the 
 loss of this the gases which escape from the chamber are 
 passed through a tower where they are brought into contact 
 with concentrated sulfuric acid running slowly over pieces 
 of coke. The sulfuric acid absorbs the nitrogen dioxide 
 
 1 This should not be learned by rote but should be written on the follow- 
 ing principles: (1) 2HNO 3 will give 2NO + H 2 O + 3O. (2) Three atoms 
 of oxygen will oxidize three molecules of sulfur dioxide to the trioxide. 
 The rest follows from this beginning. 
 
CONCENTRATION OF SULFURIC ACID 
 
 123 
 
 but the gas can be recovered by diluting the acid and bring- 
 ing it into contact with the sulfur dioxide as it comes from 
 the pyrite burners. The tower in which the nitrogen dioxide 
 is absorbed is called the Gay-Lussac tower and that in 
 which it is taken up again from the diluted acid is called the 
 Glover tower. The general arrangement of the whole 
 apparatus is shown in Fig. 29. 
 
 Leaden Chambers 
 FIG. 29. 
 
 Gay-Lussac 
 Tower 
 
 Theoretically, a small quantity if nitric acid should con- 
 vert an unlimited amount of sulfur dioxide into sulfuric 
 acid. Practically, there are unavoidable losses of the 
 oxides of nitrogen and in a well-conducted establishment 
 from 25 to 40 pounds of sodium nitrate, NaN0 3 , are re- 
 quired to furnish the nitric acid necessary for the manufac- 
 ture of a ton of sulfuric acid. 
 
 Concentration of Sulfuric Acid. Properties. The acid 
 produced by the chamber process contains 30 to 40 per 
 cent of water. For shipment and for many purposes for 
 which the acid is used a product which is nearly free from 
 water is preferred. When the chamber acid is heated, 
 water escapes at first with very little acid and it is possible 
 to carry the acid up to a strength of 96 to 98 per cent by sim- 
 ple evaporation. The evaporation is carried out at first in 
 
124 SULFUR, SELENIUM AND TELLURIUM 
 
 lead pans which are scarcely attacked by the dilute acid, 
 but as the acid becomes more concentrated the lead begins 
 to dissolve and the further concentration up to 93 to 95 
 per cent is carried out in glass, platinum or iron. This 
 gives the concentrated commercial acid, which is usually 
 slightly brown from the effect of dust and organic matter 
 and which contains arsenic, lead, iron and other impurities 
 that do not interfere with its use for most purposes. A pure 
 and stronger acid is obtained by distillation. The concen- 
 trated acid boils at 338 and contains only about two per 
 cent of water. 
 
 Considerable heat is generated when water is mixed with 
 concentrated sulfuric acid. The concentrated acid has 
 a specific gravity of L84 and it should always be diluted 
 by pouring ihe acid into water and not vice versa. If water 
 is poured upon the acid so much heat may be produced at 
 the junction of the two liquids as to create an explosion by 
 the steam which is generated. 
 
 The affinity of sulfuric acid for water is so great that it 
 abstracts the elements of water from such substances as 
 sugar or wood, charring them. 
 
 Concentrated sulfuric acid produces very painful wounds. 
 In cases of accident the acid should be removed as quickly 
 as possible by the copious application of water and after- 
 ward moist sodium bicarbonate may be applied to the 
 wound. 
 
 Because of its high boiling point and low price, sulfuric 
 acid is used to prepare more volatile acids from their salts. 
 Especially it is used in the preparation of hydrochloric and 
 nitric acids. In the reversible reactions represented by 
 the equations: 
 
 NaCl + H 2 SO 4 <=* NaHSO 4 + HC1 
 and 
 
 NaNO 3 + H 2 SO 4 + NaHSO 4 + HN0 3 
 
SULFATES. SODIUM THIOSULFATE 125 
 
 the fact that hydrochloric acid is a gas and nitric acid is 
 volatile at a moderate temperature carries the reaction to 
 the right as the hydrochloric acid or the nitric acid escapes 
 and the operation can be made practically quantitative in 
 that direction. 
 
 Large quantities of sulfuric acid are used in the treatment 
 of the mineral phosphates which are used as fertilizers. The 
 treatment makes the phosphates more soluble and so mo're 
 quickly available for the growth of crops. Sulfuric acid 
 is also used in making nitroglycerine for dynamite, in 
 making cellulose nitrate or "gun cotton, " which is the basis 
 of smokeless powder, in purifying petroleum for the manu- 
 facture of kerosene and gasolene, and in many other ways. 
 
 Sulfates and Acid Sulfates. Sulfuric acid is dibasic and 
 it forms both normal and acid salts. Thus normal sodium 
 sulfate is Na 2 SO 4 and the acid sulfate is NaHSO 4 . The 
 acid salts are strongly acid in reaction and as acids will 
 neutralize bases. They are both strong acids and salts. 
 The normal sulfates of strong bases, such as sodium sul- 
 fate, potassium sulfate and calcium sulfate, are neutral^ in 
 reaction. 
 
 Nearly all sulfates and nearly all salts of othep strong 
 acids, that is, of acids which are largely ionized/ in dilute 
 solutions, are soluble in water. The most commox^insoluble 
 sulfates are lead sulfate, PbSO 4 , barium sulfate, BaSO 4 , 
 and calcium sulfate, CaSO4. The last is sufficiently soluble, 
 however, to be a troublesome constituent in hard waters. 
 Barium sulfate, BaS0 4 , requires approximately 400,000 
 parts of water for its solution and because of this extreme 
 insolubility it is used to detect the presence of sulfuric 
 acid or a sulfate in solutions and to determine the quantity 
 of the sulfate group, SO 4 , in salts or mixtures. 
 
 Sodium Thiosulfate. A solution of sodium sulfite slowly 
 absorbs oxygen from the air and the salt is changed to 
 sodium sulfate: 
 
 Na 2 S0 3 + O = Na 2 SO 4 
 
126 SULFUR, SELENIUM AND TELLURIUM 
 
 The solution of sodium sulfite will dissolve sulfur, which 
 adds itself very much as the dxygen does, forming a thio- 
 sulf ate : 
 
 Na 2 SO 3 + S = Na 2 S 2 3 
 
 Sodium 
 thiosulfate 
 
 The prefix ihio (from Greek Oeiov, sulfur) is used to indi- 
 cae that in this compound we have a sulfate in which one 
 atom of sulfur has taken the place of an atom of oxygen. 
 
 Sodium thiosulfate crystallizes with water, the salt hav- 
 ing the composition Na2S 2 O 3 .5H 2 O. It is much used in 
 photography to "fix" pictures by dissolving that portion 
 of the silver chloride or bromide which has not been affected 
 by expos are to the light. The salt has been known for 
 a long time and when it was first discovered the acid from 
 which it was derived was called hyposulfurous acid and the 
 salt sodium hyposulfite because the acid and salt contain 
 less oxygen than sulfurous acid and sodium sulfite, in pro- 
 portion to the sulfur present in the compounds. While 
 chemists generally use the other names, the name sodium 
 I yposulfite is commonly used by pharmacists and photog- 
 raphers. 
 
 Selenium and Tellurium. On referring to the table 
 on p. 100 it will be seen that these two elements correspond 
 in their atomic weights to bromine and iodine of the halogen 
 family. They form compounds similar to the compounds 
 of sulfur, among which may be mentioned the compounds 
 having the formulas H 2 Se, H 2 Te, Se0 2 , TeO 2 , H 2 Se0 3 , 
 H 2 Te0 3 , H 2 Se0 4 , H 2 Te0 4 . 
 
 Selenium occurs in several allotropic forms, as sulfur 
 does. It is more metallic in its properties than sulfur and 
 the electrical conductivity of a film of the element is very 
 sensitive to the action of light. This property has been 
 used in an instrument for measuring the intensity of light 
 from very distant sources and has been applied in the study 
 of variable stars. 
 

 SUMMARY. SULFUR 127 
 
 Tellurium is still more metallic and in many of its proper- 
 ties it should be classed with the metals. It is sometimes 
 found associated with gold, telluride gold ores being found 
 especially in Colorado. 
 
 SUMMARY 
 
 Sulfur is found free in nature, also combined in the sulfides, 
 iron pyrites, galena and sphalerite and the sulfates, gypsum 
 and barium sulfate. 
 
 Sulfur is used in making sulfuric acid, in "sulfuring" 
 fruit for drying, in the "lime-sulfur" wash, in black gun 
 powder, in the manufacture of india-rubber and in making 
 some dyes. 
 
 Sulfur has three allotropic solid forms, rhombic and 
 needle-shaped crystals and an amorphous form. 
 
 There are two liquid forms, corresponding to the last 
 two solid forms. 
 
 Sulfur boils at 445, 
 
 Hydrogen sulfide is found in sulfur waters and in decayed 
 eggs. 
 
 For purposes of anaylsis metals are separated by the 
 conduct of solutions of their salts toward hydrogen sulfide. 
 Metals of Group I are precipitated by hydrogen sulfide 
 from acid solutions, those of Group II from alkaline solu- 
 tions, and those of Group III are precipitated from neither 
 acid nor alkaline solutions. 
 
 The formation of a precipitate is a reversible reaction 
 which goes nearly to completion because the precipitate is 
 removed from the sphere of action. 
 
 Hydrosulfides are formed by the 1 action of hydrogen sul- 
 fide on hydroxides. 
 
 Sulfur dioxide is prepared by burning sulfur or by the 
 action of concentrated sulfuric acid on copper. It is used 
 in bleaching, as a germicide and in the manufacture of 
 sulfuric acid. 
 
128 SULFUR, SELENIUM AND TELLURIUM 
 
 Sulfurous acid is formed when sulfur dioxide is dissolved 
 in water. It is a bibasic acid, giving both acid and normal 
 salts. 
 
 Sulfur trioxide is a liquid which easily polymerizes to a 
 solid form. It combines with water to form sulfuric 
 acid. 
 
 In accordance with the principle of van't Hoff-Le Chate- 
 lier a rise in temperature is unfavorable to that direc- 
 tion of a reversible reaction in which heat is evolved. 
 
 The chamber process for sulfuric acid is carried out in 
 leaden chambers into which are introduced sulfur dioxide, 
 air, nitric acid and water. 
 
 Nitric acid or nitrogen dioxide oxidizes sulfur dioxide and 
 water to sulfuric acid and the nitric oxide formed combines 
 with oxygen, acting in this way as a catalytic agent to con- 
 vert sulfur dioxide into sulfuric acid. 
 
 Sulfuric acid is concentrated in glass, platinum or iron. 
 It has a strong affinity for water. 
 
 Sulfuric acid is used in preparing hydrochloric and nitric 
 acids, in making smokeless powder, in purifying petroleum 
 and in many other ways. 
 
 Sulfuric acid is bibasic. The sulfates of calcium, lead 
 and barium are only slightly soluble. 
 
 Sodium thiosulfate is prepared by dissolving sulfur in a 
 solution of sodium sulfite. It is used to "fix" pictures in 
 photography. 
 
 Selenium and tellurium are elements of higher atomic 
 weights having properties similar to those of sulfur. 
 
 EXERCISES 
 
 1. A liter of hydrogen sulfide weighs approximately 1.52 grams. 
 How much ferrous sulfide and how much sulfuric acid will be 
 required Jbo furnish 22.4 cc. of the gas? 
 
 2. How many pounds of sulfur will be required to make a ton of 
 sulfuric acid of 70 per cent strength? How much to make a ton 
 of 95 per cent acid? 
 
EXERCISES. SULFUR 
 
 3. How many pounds of iron pyrites will be required to make a 
 ton of pure sulfuric acid? 
 
 4. How much sulfuric acid will be required to give a ton of pure 
 nitric acid, HNOs? 
 
 5. If 25 pounds of sodium nitrate are used in making a ton of 
 95 per cent sulfuric acid, how many times does each atom of nitro- 
 gen react in the process? 
 
 6. Heat is evolved when chlorine is liberated by the action of 
 oxygen or air on hydrochloric acid (p. 82). Will an increase in 
 temperature favor the formation of chlorine or the reverse? 
 
 7. Name the compounds whose formulas are given in the para- 
 graph on selenium and tellurium. 
 
CHAPTER XII 
 
 SELECTION OF MOLECULAR AND ATOMIC WEIGHTS 
 
 Combination of Gases by Volume. Two volumes of 
 hydrogen unite with one volume of oxygen to form water. 
 If the water formed is kept at the same temperature as the 
 oxygen and hydrogen which combine and this temperature 
 is high enough so that the water remains as a gas, that is, in 
 the form of steam, two volumes of steam will result from 
 the combination. One volume of hydrogen combines with 
 one volume of chlorine to form two volumes of hydro- 
 chloric acid, HC1. One volume of nitrogen unites with 
 three volumes of hydrogen to form ammonia, NH 3 . One 
 volume of chlorine unites with two volumes of oxygen to 
 form two volumes of chlorine dioxide, C1O 2 . These rela- 
 tions can be represented by the following diagrams: 
 
 2 grams 
 Hydrogen 
 
 + 
 
 71 grams 
 Chlorine 
 
 
 
 2 grams 
 Hydrogen 
 
 32 grams 
 Oxygen 
 
 2 grams 
 Hydrogen 
 
 
 36.5 grams 
 Hydrochlor- 
 ic acid 
 
 36.5 grams 
 Hydrochlor- 
 ic acid 
 
 18 grams 
 Steam 
 
 18 grams 
 Steam 
 
 130 
 
GAY LUSSAC'S LAW 
 
 131 
 
 2 grams 
 Hydrogen 
 
 + 
 
 
 2 grams 
 Hydrogen 
 
 28 grams 
 Nitrogen 
 
 2 grams 
 Hydrogen 
 
 
 
 32 grams 
 Oxygen 
 
 71 grams 
 Chlorine 
 
 32 grams 
 Oxygen 
 
 
 17 grams 
 Ammonia 
 
 17 grams 
 Ammonia 
 
 67.5 grams 
 
 Chlorine 
 
 dioxide 
 
 67.5 grams 
 
 Chlorine 
 
 dioxide 
 
 Gay Lussac's Law. 4 study of the combination of gases 
 with each other led to the discovery of Gay Lussac's Law 
 of Combination by Volume. There is always a simple 
 ratio between the volumes of gases which combine and there is 
 also a simple ratio between these volumes and the volume of 
 the product if that is a gas. 
 
 Relations between Combining Volumes and Atomic 
 Weights. The volumes for the diagrams have been selected 
 in such a manner that each volume of a gas containing 
 hydrogen contains either one gram or a whole number of 
 grams of that element. If the unit volume is selected in 
 this way it follows: 
 
132 MOLECULAR AND ATOMIC WEIGHTS 
 
 1. The ratio of the volumes is always one of simple whole 
 numbers (Gay Lussac's law). 
 
 2. The composition of every compound must be expressed 
 in terms of the atomic weights of the elements multiplied by 
 small numbers (law of combining proportions, p. 48). 
 
 3. This volume must either contain one gram atom or a whole 
 number of gram atoms, or a simple fraction of a gram atom 
 of every other element present. 
 
 As hydrogen is the unit 1 for atomic weights it is convenient 
 to select as our unit for the combining of gases that volume 
 of some compound containing one atom of hydrogen in the 
 molecule, which contains one gram atom of hydrogen. We 
 shall be likely to discover such a compound if we examine a 
 large number of compounds of hydrogen and select the one 
 which contains the smallest weight of hydrogen in a given 
 volume. Among the compounds chosen for illustration 
 hydrochloric acid contains the smallest amount of hydrogen 
 in the unit volume. An examination of many hundreds of 
 other compounds of hydrogen has revealed none which con- 
 tains less than one gram of hydrogen in this volume. We 
 may feel reasonably certain, therefore, that hydrochloric 
 acid has only one atom of hydrogen in its molecule and 
 the volume of hydrochloric acid which contains one gram 
 atom of hydrogen may be selected as our unit volume. This 
 volume is 22.4 liters (36.5 divided by 1.6398, the weight 
 of one liter of hydrochloric acid gas at and 760 mm. 
 = 22.3). 2 
 
 Having selected 22.4 liters as our unit volume it follows, 
 as has been shown, that this volume of every other gaseous 
 compound of hydrogen must contain one, two, three or 
 some other whole number of grams of hydrogen and that 
 
 1 In round numbers. O = 16 is the real basis (see p. 51). 
 
 2 Oxygen, nitric oxide, carbon dioxide and other gases which are not so 
 easily condensed to a liquid as hydrochloric acid give, in the same manner, 
 values very close to 22.4. These gases obey the gas laws more closely 
 and are considered "perfect" gases. 
 
GRAM MOLECULAR VOLUME 133 
 
 it must also contain one, two, three or some other whole 
 number of gram atoms of each element with which the 
 hydrogen is combined. This is illustrated in the diagrams: 
 22.4 liters of steam 1 contain 2 grams of hydrogen and 16 
 grams of oxygen; 22.4 liters of ammonia contain 3 grams 
 of hydrogen and 14 grams of nitrogen. 
 
 The law of combining volumes enables us to extend the 
 principle to compounds which do not contain hydrogen, 
 also. Thus 22.4 liters of chlorine dioxide contain 2 gram- 
 atoms, or 71 grams, of chlorine and 1 gram atom, or 16 
 grams, of oxygen. 
 
 As 22.4 liters has been selected as the unit volume because 
 no gaseous compound of hydrogen contains less than one 
 gram of the element in that volume, so we may select for 
 other elements the smallest weight of the element found in 
 22.4 liters of the various gaseous compounds of the element. 
 If the compounds selected in this manner contain only one 
 atom of the element in the molecule, the weight of the 
 element in the unit volume must be one gram atom, and 
 the atomic weight of the element follows from this. Thus 
 there is no gaseous compound of oxygen which contains 
 less than 16 grams of the element in 22.4 liters. The atomic 
 weight of oxygen is selected as 16, therefore, and in a 
 similar manner 35.5 has been chosen as the atomic weight 
 of chlorine, 14 as the atomic weight of nitrogen and 32 as 
 the atomic weight of sulfur. Other atomic weights have 
 been chosen in the same way. 
 
 The fact that the atomic weights chosen by the principles 
 which have been outlined have led to an orderly arrange- 
 ment of the elements (the Periodic System, p. 162) in which 
 there has been shown to be a close relation between the 
 atomic weights and the properties of the elements is a very 
 
 1 By this is meant, of course, that volume of steam which would fill 
 22.4 liters if it could be brought to normal conditions (0 and 760 mm.) 
 without condensation to a liquid. 
 
134 MOLECULAR AND ATOMIC WEIGHTS 
 
 strong reason for believing that these principles are correct 
 and that the numbers chosen actually represent the rela- 
 tive weights of the atoms of the elements. 
 
 Gram Molecular Volume. The molecular wei-ght of a 
 compound is the sum of the atomic weights of the elements 
 of which it is composed, each atomic weight being taken as 
 many times as there are atoms of the element. Thus the 
 molecular weight of water, H 2 O, is 2 + 16 = 18; the mole- 
 cular weight of ammonia, NH 3 , is 3 + 14 = 17; the mole- 
 cular weight of chlorine dioxide, C1O 2 , is 35.5 + (2 X 16) 
 = 67.5; the molecular weight of sulfuric acid, H 2 SO 4 , 
 is (2 X 1) + 32 + (4 X 16) = 98. A gram molecule of a 
 compound weighs as many grams as the molecular weight. 
 A gram molecule of water weighs 18 grams. A gram mole- 
 cule of sulfuric acid weighs 98 grams. 
 
 It has been shown in the preceding paragraph that the 
 unit volume,* 22.4 liters, has been so chosen that it must 
 contain as manjr gram atoms of hydrogen and of each of 
 the other elements as there are atoms of the element in 
 the compound. Thus 22.4 liters of steam contain (2X1) 
 + 16 = 18 grams of steam and 22.4 liters of chlorine dioxide 
 contain 35.5 + (2 X 16) = 67.5 grams of the gas. It 
 follows, then, that 22.4 liters of any gas contain one gram 
 molecule of the gas. Because of this relation, 22.4 liters is 
 called a gram-molecular volume. If we can determine the 
 weight of 22.4 liters of any gas under standard conditions 1 
 we have the weight of a gram molecule and so the molecular 
 weight of the compound or element. Conversely, if we 
 know the formula of any gaseous compound we have only 
 to add the weights of the atoms together to find the 
 weight in grams of 22.4 liters of the compound. 
 
 An additional meaning which may be given to formulas 
 
 1 At and 760 mm. If the determination is made at any other tem- 
 perature or pressure, the result may be calculated for standard conditions 
 by means of the gas laws (p. 37). 
 
AVOGADRO'S LAW 135 
 
 of gaseous compounds follows from the foregoing. Just 
 as a formula may represent a gram molecule of a compound 
 (p. 56) it may also represent a gram^molecular volume. 
 Thus the equation: 
 
 N 2 + 3H 2 = 2NH ? 
 
 Nitrogen Hydrogen Ammonia 
 
 not only means that one molecule of nitrogen, N 2 (see 
 below), combines with three molecules of hydrogen, H 2 , to 
 form two molecules of ammonia, NH 3 , but it also means that 
 a gram-molecular volume (22.4 liters) of nitrogen combines 
 with three gram-molecular volumes (3 X 22.4 liters) of 
 hydrogen to form two gram-molecular volumes (2 X 22.4 
 liters) of ammonia. 
 
 Number of Molecules in the Unit Volume of a Gas. 
 Avogadro's Law. If 22.4 liters of any gas contain a gram 
 molecule of the gas it must follow that if we could decrease 
 the unit volume till it contained only 1000 molecules of the 
 gas the same volume would contain 1000 molecules of any 
 other gas, for the weights of this small volume of different 
 gases must still be proportional to their molecular weights. 
 In other words, equal volumes of all gases under the same 
 conditions of temperature and pressure contain the same num- 
 ber of molecules. This is known as Avogadro's law because 
 it was discovered by an Italian chemist, Avogadro, in 1811. 
 He based his hypothesis on the law of combining volumes. 
 It has been pointed out (p. 33) that the diffusion of gases 
 and other phenomena which have led to the kinetic theory 
 of gases are best explained on the supposition that equal 
 volumes contain equal numbers of molecules. So many 
 different phenomena are consistently explained on this 
 basis that the law is now considered to be thoroughly 
 established. 
 
 The actual number of molecules in a given molecular 
 volume must be, from the figures given on p. 33, 22.4 X 
 1000 X 2.71 X 10 19 (or 607,000,000,00Q>QQQ,OQp,000,000).. 
 
136 MOLECULAR AND ATOMIC WEIGHTS 
 
 Formulas of Oxygen and of Other Elementary Gases. 
 The gram-molecular volume of oxygen weighs 32 grams. 
 From this the formula of oxygen must be O 2 . In a similar 
 manner the formulas of hydrogen, chlorine and nitrogen are 
 H 2 , C1 2 and N 2 . 
 
 The same conclusions may be reached as follows on the 
 basis of Avogadro's law: 
 
 Suppose that a volume of hydrogen containing 1000 mole- 
 cules be taken. It will combine with the same volume of 
 chlorine, which must, according to the law, also contain 
 1000 molecules. The combination will give two volumes of 
 hydrochloric acid, which must contain 2000 molecules. 
 In other words, one molecule of hydrogen with one molecule 
 of chlorine gives two molecules of hydrochloric acid. 
 
 1000 + 1000 = 1000 + 1000 
 
 Molecules Molecules Molecules Molecules 
 
 of of of of 
 
 hydrogen chlorine hydrochloric hydrochloric 
 acid acid 
 
 Since each molecule of hydrochloric acid must contain 
 at least one atom of hydrogen and one atom of chlorine, 
 each molecule of hydrogen and each molecule of chlorine 
 must contain at least two atoms of the element. 
 
 SUMMARY 
 
 The volumes of gases which react always bear a simple 
 ratio to each other and the ratios between these volumes 
 and the volume of the product, if that is a gas, is also 
 simple (Gay Lussac's Law). 
 
 Twenty-two and four-tenths liters of any gaseous 
 compound under standard conditions contain one gram 
 molecule of the compound and this volume is called a 
 gram-molecular volume. 
 
 The smallest number of grams of an element found in any 
 
EXERCISES. MOLECULAR WEIGHTS 137 
 
 of its gaseous compounds is chosen as the weight of a 
 gram atom of the element and this fixes the atomic weight. 1 
 The formulas of oxygen, hydrogen, chlorine and nitrogen 
 are O 2 , H 2 , C1 2 , N 2 . 
 
 EXERCISES 
 
 1. What is the weight of a gram-molecular volume of the 
 following gases: Chlorine monoxide, C1 2 0, hydrobromic acid, 
 hydrogen sulfide, sulfur dioxide, sulfur trioxide, hydrogen tellu- 
 ride? 
 
 2. A gram-molecular volume of air weighs about 29 grams. 
 (Notice the relation of this to the weight of a gram-molecular 
 volume of nitrogen, 28 grams, and of oxygen, 32 grams.) What is 
 the weight of each of the gases mentioned as compared with the 
 weight of the same volume of air? 
 
 & The weight of a liter of ozone is approximately 2.15 grams; 
 what is the formula? 
 
 4. At 200 the weight of that volume of iodine vapor which 
 would fill one liter if it could be cooled to under a pressure of 
 760 mm. is approximately 11.3 grams. What is the formula of 
 iodine at that temperature? 
 
 5. At 1000 the weight of that volume of iodine vapor which 
 would occupy one liter under standard conditions is 5.7 grams. 
 What is the formula of iodine at that temperature? 
 
 6. What weight of sulfur will be required to give 22.4 liters of 
 sulfur dioxide? What volume of oxygen? 
 
 7. What volume of oxygen will be required to give 22.4 
 liters of sulfur trioxide, if it could exist as a gas under standard 
 conditions? 
 
 1 Approximately. Accurate atomic weights are determined by the 
 quantitative analysis of compounds of the element, especially of com- 
 pounds with chlorine, bromine or iodine. A few elements have no gaseous 
 compounds, or have none containing a single atom of the element in a 
 molecule. The atomic weights of these elements are fixed by the law of 
 Dulong and Petit (p. 266). 
 
CHAPTER XIII 
 GROUP V: NITROGEN 
 
 Occurrence of Nitrogen. Nitrogen forms nearly four- 
 iifths of the volume of air and a little more than three- 
 fourths of its weight. As the pressure of the air at sea 
 level is nearly 15 pounds to the square inch, the weight of 
 nitrogen above one square inch is about 11 pounds, or more 
 than three-fourths of a ton over one square foot and more 
 than 20,000,000 tons over one square mile. In spite of this 
 enormous quantity of the element in the atmosphere it is so 
 difficult to cause nitrogen to combine with other elements 
 that compounds of nitrogen are comparatively expensive. 
 
 Nitrogen is an essential constituent of all living bodies, 
 whether animal or vegetable. Corn, wheat, oats, grass and 
 most vegetables must secure nitrogen for their growth 
 from compounds of nitrogen present in the soil, especially 
 from potassium nitrate, KNO 3 , calcium nitrate, Ca(NO 3 )2, 
 or ammonia, NH 3 . Apparently none of the cereals can 
 utilize nitrogen of the air directly and they will not grow 
 on a soil that does not contain compounds of nitrogen avail- 
 able for their use. 
 
 Alfalfa, clover and some other legumes may, however, 
 utilize the nitrogen of the air with the aid of bacteria which 
 grow on their roots, producing nodules. A small amount of 
 -oxygen and nitrogen combine during every lightning flash 
 and the oxides of nitrogen formed combine with water and 
 .ammonia in the air to form ammonium nitrite, NH 4 NO2. 
 This is carried down by rain and becomes available for the 
 .growth of plants. 
 
 Organic matter in the soil decays more or less rapidly 
 
 138 
 
NATURAL HISTORY OF NITROGEN 139 1 
 
 under the influence of bacteria. If air is present, potassium 
 nitrate, KNO 3 , or calcium nitrate, Ca(N0 3 ) 2 , is formed. 
 In the absence of air, where the conditions favor reduction 
 rather than oxidation, ammonia, NH 3 , is formed. Either 
 of these may then be available for the growth of new mate- 
 rial. The course of nitrogen in nature which has been out- 
 lined may be seen from the following diagram: 
 
 Leguminous plants with 
 
 the help of bacteria 
 Atmospheric 
 electricity Plants 
 
 Atmospheric 
 Nitrogen 
 
 Nitrates 
 
 Organic compounds 
 of Nitrogen 
 
 Denitrifying NiTtrifying ' Decayjof plant" 
 
 bac teria and ani mal tissues 
 
 or dis I tillation 
 
 | Nitrifying 
 
 Nitrites U 
 
 bacteria 
 
 Preparation and Properties of Nitrogen. When phos- 
 phorus is burned in a confined portion of air the oxygen is 
 removed and nearly pure nitrogen remains. The oxygen 
 may also be completely removed from air by passing it 
 over heated metallic copper. Nitrogen is colorless and 
 odorless and will not support combustion. It is very 
 much more inert than the elementary gases and other 
 elements which have been studied and this is evidently 
 one reason why such large quantities of nitrogen are found 
 uncombined in the atmosphere. 
 
 At the very high temperatures produced by electric 
 sparks a portion of the nitrogen and oxygen of the air 
 through which the sparks are passed combines to form 
 nitric oxide, NO: 
 
 N 2 + O 2 = 2NO 
 Nitrogen and hydrogen will combine at moderate tem- 
 
140 NITROGEN 
 
 peratures with the aid of metallic osmium or uranium as a 
 catalyzer : 
 
 N 2 + 3H 2 = 2NH 3 
 
 Both of these processes are now in technical use and 
 promise to be increasingly important. 
 
 Ammonia. When organic matter, such as stable manure, 
 or sewage, decomposes with exclusion of air, through the 
 action of bacteria, a part of the nitrogen is converted into 
 ammonia, NH 3 . Ammonia is also formed when almost any 
 kind of natural organic matter containing nitrogen is heated 
 to a high temperature. In this manner ammonia is formed 
 when coal is heated for the manufacture of illuminating gas 
 or of coke and nearly all of the ammonia and ammonium 
 salts of commerce come from this source. 
 
 Ammonia is most readily prepared for laboratory pur- 
 poses by warming a concentrated solution of the gas in 
 water. The gas may be dried by passing it through a tube 
 filled with solid potassium hydroxide or sodium hydroxide, 
 or with quicklime. 
 
 Ammonia is a colorless gas with a pungent odor. Is the 
 gas heavier or lighter than air, and in what proportion? 
 A gram-molecular volume of air weighs about 29 grams. 
 
 Water at the freezing point will dissolve 1000 times its 
 volume of ammonia. The solution is lighter than water 
 and gives off ammonia very readily when it is warmed. 
 The solution of ammonia reacts alkaline toward litmus and 
 neutralizes acids. These properties and others have led 
 to the view that the ammonia combines with the water, in 
 part, to form ammonium hydroxide, NH 4 OH : 
 
 NH 3 + HOH = NH 4 OH 
 or, graphically, 
 
 H \ H \ /^ 
 
 H^N + H O H = H-^N(^ 
 
 W W X H 
 
AMMONIUM SALTS 14 1 
 
 The nitrogen, which is trivalent in ammonia, becomes 
 quinquivalent in ammonium hydroxide. The ammonium 
 hydroxide gives in a solution the positive ammonium ion, 
 NH 4 +, and the negative hydroxide ion, OH~. It is the 
 hydroxide ion, of course, which causes the alkaline reaction 
 of the solution (p. 74). 
 
 Ammonium Salts. Ammonia combines directly with 
 acids to form ammonium salts in which the univalent 
 ammonium group, NH 4 , takes the place usually occupied 
 by a metal: 
 
 NH 3 + HC1 = NH 4 C1 
 
 Ammonium chloride 
 
 Ammonium chloride resembles common salt, NaCl, in 
 taste and in many of its properties. How would it react 
 with concentrated sulfuric acid? 
 
 2NH 3 + H 2 S0 4 = (NH 4 ) 2 SO 4 
 
 Ammonium sulfate 
 
 NH 3 + HN0 3 = NH 4 N0 3 
 
 Ammonium nitrate 
 
 Or, graphically, 
 
 H \ H \ / H 
 
 H^N + H-C1 = H^N( 
 
 W W X C1 
 
 In each case one molecule of ammonia combines with one 
 atom of hydrogen from the acid. For the formation of a 
 normal salt a monobasic acid requires one molecule of 
 ammonia, a bibasic acid two molecules and a tribasic acid 
 three molecules. 
 
 Action of Bases on Ammonium Salts. When an ammo- 
 nium salt is treated with a base the usual interaction with 
 partial exchanges of ions takes place: 
 
 NH 4 C1 + NaOH <= NaCl + NH 4 OH 
 If little water is present, the ammonium hydroxide 
 decomposes and the ammonia escapes: 
 NH 4 OH * NH 3 + HO 
 
142 NITROGEN 
 
 This is the usual commercial method of preparing 
 ammonia, but calcium hydroxide, Ca(OH) 2 , is used instead 
 of sodium hydroxide. Why? 
 
 Synthesis of Ammonia. It has been known for a long 
 time that when electric sparks are passed through a mixture 
 of nitrogen and hydrogen a minute quantity of ammonia 
 is formed: 
 
 N 2 + 3H 2 <= 2NH 3 
 
 But ammonia gas is also decomposed into nitrogen and 
 hydrogen by the passage of electric sparks and the equilib- 
 rium in this reversible reaction is so very far on the side 
 toward decomposition that it is wholly impracticable to 
 prepare ammonia by this method. 
 
 A practical commercial synthesis was finally discovered 
 by a careful study and application of the following facts 
 and principles: 
 
 First.- The reaction is exothermic, that is, heat is gen- 
 erated when nitrogen and hydrogen combine. 
 
 Second. Because heat is generated in the combination, 
 an increase in the temperature favors the decomposition 
 &iid not the formation of ammonia (principle of Van't 
 Hoff-Le Chatelier, p. 120). 
 
 Third. The combination of the gases takes place very 
 slowly at any temperature at which the equilibrium would 
 be favorable; that is, at any temperature at which the 
 combination goes far enough to be commercially profitable. 
 
 Fourth. Osmium, uranium and some other metals 
 catalyse the reaction (p. 9) and make it practicable at a 
 temperature of 400 or below. 
 
 Fifth. Four volumes of the mixture of nitrogen and 
 hydrogen give only two volumes of ammonia. An increase 
 of the pressure, therefore, favors the combination (prin- 
 ciple of Van't Hoff-Le Chatelier). 
 
 On the basis of .these facts and principles the commercial 
 
OXIDATION OF AMMONIA 143 
 
 manufacture of ammonia is carried out by passing a highly 
 compressed mixture of nitrogen and hydrogen over osmium, 
 uranium or some other catalyzer. A commission of the 
 Government of the United States has recommended the 
 use of this process on a large scale for the manufacture of 
 ammonia, which is then to be oxidized to nitric acid (see 
 below). There is strong reason for believing that this and 
 other synthetic processes for making ammonia and nitric 
 acid were extensively used in Germany during the Great 
 War and that it is pnly through the use of these processes 
 that her speedy defeat, through lack of munitions, was 
 prevented. 
 
 Oxidation of Ammonia to Nitric Oxide and Nitric Acid. 
 Smokeless powder, gun cotton, nitroglycerine, trinitro- 
 toluene ("T. N. T. ") gunpowder and all other explosives, 
 used in blasting or in warfare, except some of the primers, 
 and "ammonal," a mixture of ammonium chloride and 
 aluminum powder, require nitric acid or a nitrate for their 
 manufacture. Before the nineteenth century the world's 
 supply of nitric acid and nitrates was obtained from salt- 
 peter, KNO 3 , and a few other natural nitrates. For many 
 years before the Great War the principal source of supply 
 was Chili saltpeter, NaNOa. This natural supply will be 
 exhausted within a comparatively few years and chemists 
 have been searching eagerly for methods of producing nitric 
 acid from other materials. One of the most promising of 
 the methods discovered was developed by Ostwald. It 
 consists in passing a mixture of ammonia and air over 
 heated platinized asbestos, which serves as a catalyzer. 
 The hydrogen of the ammonia is burned to water and the 
 nitrogen to nitric oxide. The latter can be converted to 
 nitric acid without much difficulty (p. 151) : 
 4NH 3 + 5O 2 = 6H 2 + 4NO 
 
 About 90 per cent of the nitrogen of the ammonia may be 
 converted into nitric acid by this process. 
 
144 NITROGEN 
 
 Nitric Acid. When organic matter decays under the 
 influence of bacteria in a soil to which air has free access 
 a part of the nitrogen is converted into nitrates. The 
 nitrates most common in the soil are those of potassium, 
 sodium and calcium, KNOs, NaNO 3 and Ca(NO 3 ) 2 . 
 
 In an arid region in Chili and Peru, in South America, 
 enormous beds of sodium nitrate, NaNO 3 , have been formed 
 and for many years this mineral has been the almost ex- 
 clusive source from which nitric acid and the saltpeter or 
 potassium nitrate, KNO 3 , used for gunpowder, have been 
 manufactured (see above). 
 
 Nitric acid is manufactured by mixing sodium nitrate 
 with concentrated sulf uric acid and distilling the mixture : 
 
 NaN0 3 + H 2 S0 4 = NaHSO 4 + HNO 3 
 
 Acid sodium 
 sulfate 
 
 Nitric acid is volatile at a much lower temperature than 
 sulf uric acid and, because of this, the reversible reaction 
 may be carried to completion in the direction which is 
 desired. 
 
 Properties of Nitric Acid. Pure nitric acid boils at 86. 
 This pure acid is difficult to prepare and decomposes easily, 
 especially when exposed to the light: 
 
 4HN0 3 = 4NO 2 + O 2 + 2H 2 O 
 
 Nitrogen 
 dioxide 
 
 The nitrogen dioxide gives a yellow color to nitric acid 
 containing it. 
 
 The ordinary concentrated nitric acid of the laboratory 
 contains 30 to 35 per cent of water. Such an acid boils 
 at about 120 and is very much more stable than the more 
 concentrated acid. 
 
 Oxidation with Nitric Acid. Pure nitric acid is a very 
 powerful oxidizing agent. If a little of the acid is put in a 
 test-tube and some wool, hair or feathers is placed in the 
 mouth of the tube, on boiling the acid the wool, or the other 
 
NITRIC ACID AND METALS 145 
 
 materials mentioned, will take fire and burn when the 
 vapor of the acid comes in contact with them. 
 
 Burning charcoal will continue to burn if thrust beneath 
 the surface of the pure acid. 
 
 Tin is oxidized by ordinary strong nitric acid to meta- 
 stannic acid, a white powder containing hydrogen and oxy- 
 gen as well as tin. The nitric acid is reduced to nitric oxide, 
 NO, and nitrogen dioxide, N(>2. 
 
 Action of Nitric Acid on Metals. From the action of 
 hydrochloric or sulfuric acid on zinc and iron we should 
 expect that nitric acid would give, with these metals, nitrates 
 and hydrogen. The action of sulfuric acid on copper, 
 however, has shown us that the action of an acid on a metal 
 may take a different course. In that case the sulfuric 
 acid is reduced and no hydrogen is liberated but sulfur 
 dioxide, a reduction product of sulfuric acid, instead. A 
 high temperature is required for sulfuric acid to act rapidly. 
 Nitric acid acts in a similar manner, at ordinary tempera- 
 tures, on all metals which are attacked by it. This effect 
 is easily understood when we remember that nitric acid is 
 much more vigorous than sulfuric acid, as an oxidizing 
 agent. 
 
 We may explain the action of the acid by supposing that 
 the nitric acid oxidizes the metal and is at the same time 
 reduced to one of the lower oxides of nitrogen and water, 
 and that the oxide of the metal reacts with more of the acid 
 to form a nitrate. 
 
 Or we may suppose that the metal displaces the hydrogen 
 of the acid forming a nitrate and that the hydrogen at the 
 moment of liberation acts upon more of the acid, reducing 
 it. The following equations for the action of nitric acid on 
 copper illustrate the two methods of explaining the action. 
 The formulas in brackets indicate substances which may 
 be supposed as intermediate in the reaction but for whose 
 
 formation we have no direct evidence. 
 10 
 
146 NITROGEN 
 
 First explanation : 
 
 2HNO 3 + 3Cu = [3CuO] + 2NO + H 2 O 
 6HN0 3 + [3CuQ] = 3Cu(NO 3 ) 2 + 3H 2 O 
 Adding after eliminating [3CuO]: 
 
 8HNO 3 + 3Cu = 3Cu(NO 3 ) 2 + 2NO + 4H 2 
 Second explanation: 
 
 2HNO 3 + Cu = Cu(NO 3 ) 2 + [2H] 
 2HNO 3 + [6H] = 2NO + 4H 2 O 
 
 Multiplying the first equation by 3 and adding it to the 
 second after eliminating [6H] we have : 
 
 8HN0 3 + 3Cu = 3Cu(NO 3 ) 2 + 2NO + 4H 2 O 
 
 It will be noticed that the final equation is the same by 
 one explanation as by the other and that the products of 
 the reaction give us no means of deciding which explanation 
 is true. When nitric acid acts on tin it oxidizes the tin to 
 metastannic acid and tin nitrate is not formed. On the 
 other hand, when nitric acid acts on iron the nitric acid is 
 reduced to ammonia, NH 3 . It seems that the first expla- 
 nation agrees with the action on tin better than the second, 
 while the second gives a better account of the action on iron. 
 
 Write the equations for the following reactions: iron 
 and sulfuric acid; hydrogen and nitric acid giving ammonia; 
 ammonia and sulfuric acid; ferrous sulfate and sodium 
 hydroxide giving ferrous hydroxide, Fe(OH) 2 ; ammonium 
 sulfate and sodium hydroxide. 
 
 Aqua Regia. Neither nitric nor hydrochloric acid alone 
 will dissolve gold or platinum but the metals dissolve 
 easily in a mixture of the two acids. The nitric acid as an 
 oxidizing agent liberates chlorine from the hydrochloric 
 acid and the chlorine combines with the metals forming 
 soluble chlorides. The mixture of acids is called aqua 
 regia, meaning "royal water." The name refers to the 
 designation of gold and platinum as " noble" metals. 
 
NITROUS ACID. NITROUS OXIDE 147 
 
 Nitrites. Nitrous Acid. When sodium nitrate, NaNO 3 , 
 or potassium nitrate, KN0 3 , is heated with metallic lead 
 or copper either salt is reduced to a nitrite, NaNO 2 or KNO 2 . 
 The addition of an acid to a solution of a nitrite liberates 
 nitrous acid, HNO 2 . If the solution is dilute, the nitrous 
 acid remains dissolved in the water. If the solution is 
 concentrated, the nitrous acid decomposes into nitrous an- 
 hydride, N 2 3 , and water. The nitrous anhydride, in turn, 
 decomposes into nitric oxide, NO, and nitrogen dioxide, 
 NO 2 . 
 
 Oxides of Nitrogen. There are six oxides of nitrogen 
 but two of them have the same composition and pass so 
 readily each into the other that they are often spoken of as 
 a single oxide. The oxides are: 
 
 Nitrous oxide, N 2 O 
 
 Nitric oxide, NO 
 
 Nitrous anhydride N 2 3 , 
 f Nitrogen dioxide, NO 2 
 I Nitrogen tetroxide, N 2 O 4 
 
 Nitric anhydride, N 2 O 5 
 
 All of these oxides except the last may be prepared by 
 the reduction of nitric acid. 
 
 Nitrous Oxide. When ammonium nitrate, NH 4 NO 3 , is 
 heated the hydrogen of the ammonium group, NH 4 , com- 
 bines with part of the oxygen of the nitrate group, NO 3 , 
 while the oxygen remains combined with the two nitrogen 
 atoms as nitrous oxide, N 2 0. 
 
 NH 4 NO 3 = N 2 O + 2H 2 O 
 
 Nitrous oxide is a colorless gas with a slightly sweetish 
 taste. When inhaled in small quantity it produces an 
 intoxicating effect and it is called, for this reason, laughing 
 gas. In larger quantities it produces insensibility and is 
 used for minor surgical operations, especially for the extrac- 
 tion of teeth. By use with oxygen, to prevent suffocation, 
 
148 NITROGEN 
 
 the anesthesia may be continued for a longer time and it 
 is claimed that the gas has some marked advantages over 
 ether or chloroform, which are more often used. 
 
 Nitrous oxide supports combustion and causes a glowing 
 splinter to inflame, but its effect in this regard is much less 
 vigorous than that of oxygen. 
 
 Nitric Oxide. The preparation of nitric oxide by the 
 action of nitric acid on copper has been discussed and need 
 not be described again (p. 145). 
 
 At high temperatures nitrogen and oxygen unite in 
 accordance with the reversible reaction represented by the 
 equation : 
 
 N 2 + 2 <= 2NO 
 
 The reaction is endothermic, that is, heat is absorbed as 
 the reaction proceeds. It has been pointed out that in 
 reversible reactions when heat is generated by the combina- 
 tion of elements or compounds an increase in temperature 
 shifts the equilibrium toward the decomposition of the 
 product formed (principle of van't Hoff-Le Chatelier, 
 pp. 120 and 142). The same principle leads to the con- 
 clusion that when heat is absorbed as the elements or sub- 
 stances unite an increase in the temperature causes the 
 equilibrium to shift toward the formation of the compound. 
 This conclusion has been confirmed for the combination 
 of oxygen and nitrogen as shown in the following table, 
 which gives the per cent of the oxygen and nitrogen com- 
 bining when a mixture of equal volumes of the gases is 
 heated : 
 
 Temperature Per cent of NO calculated Observed 
 1538 0.35 0.37 
 
 1922 0.98 0.98 
 
 2402 2.37 2.23 
 
 2927 4.43 About 5.00 
 
NITRIC OXIDE 149 
 
 As the combination of nitrogen and oxygen to form nitric 
 oxide is the first and most difficult step in the manufacture 
 of nitric acid from atmospheric nitrogen, the conditions 
 best suited for the reaction have been carefully studied. 
 It is evident from a consideration of the table that the best 
 arrangement will be one in which the gases are brought to 
 the highest possible temperature and then cooled quickly. 
 If cooled slowly the high percentage combination secured 
 at the high temperatures would be lost because the equili- 
 brium shifts toward the decomposition of the nitric oxide 
 into nitrogen and oxygen and at these high temperatures 
 the reaction still goes rapidly in either direction. If the gas 
 can be brought to a lower temperature without decomposi- 
 tion, the decomposition becomes so slow that the gas formed 
 at the higher temperatures is practically all saved. Even 
 at a temperature of 725 it is estimated that it would take 
 80 years for one-half of the gas corresponding to the equili- 
 brium to be formed or decomposed, while at 1825 one-half 
 of the equilibrium amount would be formed or decomposed 
 in 5 seconds an illustration of the enormous acceleration 
 of reactions at high temperatures. Nitric oxide contains 
 the same amount of oxygen in a given volume that nitrous 
 oxide, N 2 O, does, but the oxygen seems to be held in a very 
 different manner. Nitrous oxide causes a glowing splinter 
 to inflame while nitric oxide will extinguish it. Phosphorus, 
 if well ignited, will, however, burn brilliantly in nitric 
 oxide and a mixture of the vapor of carbon disulfide with 
 nitric oxide will burn with a brilliant blue flash. 
 
 Nitric oxide is a colorless gas. It combines directly with 
 oxygen to form nitrogen dioxide, NC>2, a dark brown gas 
 (see below). 
 
 Manufacture of Nitric Oxide and Nitric Acid by Means of 
 the Electric Arc. For the manufacture of nitric oxide and 
 nitric acid the high temperature necessary to secure the 
 rapid combination of nitrogen and oxygen is secured by 
 
150 NITROGEN 
 
 passing air quickly through a long, flaming, electric arc. 
 To secure the rapid cooling of the nitric oxide, which is 
 very essential in accordance with the preceding paragraph, 
 the air, after being heated in the electric flame, is carried 
 rapidly out of the flame by the action of a powerful magnetic 
 field. In Norway, where water power is cheap on account 
 of the numerous waterfalls, it has been found practicable to 
 manufacture nitric oxide and, from this, nitric acid, by 
 this method. The consumption of energy is rather large 
 in proportion to the quantity of nitric acid produced and it 
 is uncertain whether this process can compete permanently 
 with the manufacture of nitric acid by the oxidation of syn- 
 thetic ammonia (p. 143). 
 
 The formation of nitric acid is completed by the action of 
 the oxygen still remaining in the mixture and of water (see 
 below under nitrogen dioxide) . 
 
 Nitrous Anhydride. The decomposition of nitrous acid, 
 HNO 2 , into nitrous anhydride, N2O 3 , and the further de- 
 composition of nitrous anhydride into nitric oxide, NO, 
 and nitrogen dioxide, N0 2 , have been referred to under 
 nitrites and nitrous acid. Nitrous anhydride is also formed 
 as a dark green or blue liquid when nitric oxide and oxygen 
 are brought together in the right proportions at a tempera- 
 ture below 0. In what proportion, by volume, should 
 nitric oxide and oxygen be mixed to give the compound? 
 
 A decomposition like this, in which a substance decom- 
 poses and the products of the decomposition recombine 
 when the conditions are reversed, is called dissociation. 
 Nitrogen Dioxide, NO 2 , and Nitrogen Tetroxide, N 2 O 4 , 
 When nitric oxide, NO, is mixed with oxygen, the two gases 
 combine directly to form nitrogen dioxide, NO 2 . In what 
 proportion should the gases be mixed? At ordinary tem- 
 peratures nitrpgen dioxide combines with itself, in part, to 
 form the polymer, nitrogen tetroxide, N 2 O 4 . At 150 or 
 above this decomposes entirely into the dioxide. How may 
 
NITROGEN TETROXIDE 151 
 
 the true formula at a given temperature be determined? 
 At low temperatures the mixture of the dioxide and 
 tetroxide condenses to a reddish brown liquid and at 
 10.5 this freezes to colorless crystals, which seem to 
 consist entirely of the tetroxide, N 2 4 . 
 
 Nitrogen Tetroxide or Nitrogen Dioxide and Water. 
 As nitrous anhydride, N 2 Oa, dissolves in water forming 
 nitrous acid, HNO 2 , nitrogen tetroxide dissolves in cold 
 water with the formation of a mixture of nitric acid and 
 nitrous acid. What takes place in the two cases is, perhaps, 
 most easily understood by means of the following graphical 
 formulas : 
 
 Nitrous anhydride, = N O4-N = 
 
 Water, H+O - H 
 
 or, N 2 O 3 + H 2 O = 2HN0 2 
 
 ,fl 
 
 Nitrogen tetroxide, O = N OrN3 
 
 Water, H-f-OH 
 
 or, N 2 O 4 + H 2 O = HNO 2 + HNO 3 
 
 With warm water nitrogen dioxide gives nitric oxide, NO, 
 <*nd nitric acid. This may take place as follows: 
 
 Nitrogen tetroxide, O = N j 
 
 : 
 
 Water, H - O H 
 
 Nitrogen dioxide, 
 
 . 
 
 The designation anhydride means "without water" and it 
 is applied either to compounds which are formed by the 
 removal of water from an acid or to compounds which com- 
 bine with water to form acids. 
 
152 NITROGEN 
 
 SUMMARY 
 
 Nitrogen is found free, in organic compounds and in 
 nitrites and nitrates. 
 
 Legumes with the aid of bacteria can utilize the nitrogen 
 of the air. 
 
 Decomposition of organic matter by bacteria in the 
 absence of air gives ammonia. In the presence of air, as 
 in a porous soil, it gives nitrites and nitrates. 
 
 Decomposition of organic matter by heat gives ammonia. 
 
 Ammonia is obtained as a by-product in heating coal for 
 the manufacture of gas or coke. 
 
 Ammonia combines with water to form ammonium hy- 
 droxide. It combines with acids to form ammonium salts. 
 Bases liberate ammonia from such salts. 
 
 Ammonia may be prepared by the direct union of nitro- 
 gen and hydrogen. The reaction is exothermic and the 
 formation is favored by a low temperature and by high 
 pressure. A catalyzer must be used. 
 
 Ammonia may be oxidized by the oxygen of air to nitric 
 oxide, with platinized asbestos as a catalyzer. Nitric acid 
 is used in the manufacture of practically all explosives. 
 
 Until recently nitric acid and nitrates have been prepared 
 almost exclusively from Chili saltpeter. 
 
 Nitric acid is prepared by distilling a mixture of sodium 
 nitrate and sulfuric acid. 
 
 Nitric acid is a powerful oxidizing agent. It gives with 
 metals a nitrate and nitric oxide, nitrogen dioxide or ammo- 
 nia; rarely, nitrous oxide. The action may be explained as 
 an oxidation of the metal or as a reduction of the acid by 
 liberated hydrogen. 
 
 Aqua regia dissolves gold or platinum. 
 
 Sodium or potassium nitrate may be reduced to a nitrite 
 by heating with lead or copper. 
 
 Nitrites give nitrous acid with an acid. Nitrous acid 
 
EXERCISES. NITROGEN 153 
 
 decomposes to water and nitrous anhydride and the last 
 to nitric oxide and nitrogen dioxide. 
 
 There are six oxides of nitrogen. All can be prepared 
 from nitric acid. 
 
 Nitrous oxide is formed by heating ammonium nitrate. 
 It is used as an anesthetic. 
 
 Nitric oxide is prepared by the action of nitric acid on 
 copper. 
 
 Nitric oxide is manufactured by passing air through an 
 electric arc. The reaction is endothermic and is favored 
 by a very high temperature and rapid cooling of the product. 
 
 Nitrous anhydride is formed by the combination of 
 nitric oxide and oxygen at a low temperature. It disso- 
 ciates to nitric oxide and nitrogen dioxide. 
 
 Dissociation is a reversible decomposition. 
 
 Nitrogen dioxide is formed by the combination of nitric 
 oxide and oxygen. 
 
 Nitrogen dioxide gives nitrous and nitric acids with 
 cold water; nitric acid and nitric oxide with hot water. 
 
 EXERCISES 
 
 1. What are the acids which may be formed from the following 
 anhydrides: C1 2 0, C1 2 7 , I 2 5 , S0 2 , SO 3 , P 2 5 . Phosphoric 
 anhydride, t P 2 5 , forms three different acids by the addition of 
 one, two or three molecules of water. Write the equations for the 
 formation of each of these. 
 
 2. What is the weight of 22.4 liters of the mixture of nitric oxide 
 and nitrogen dioxide which would combine to form nitrous an- 
 hydride? 
 
 3. How many grams of copper will be required to prepare a 
 gram-molecular volume of nitric oxide? How many grams of 
 nitric acid, of sp. gr. 1.20 and containing 32 per cent of the pure 
 acid, will be required? 
 
 4. How many grams of copper will be required to reduce 10 
 grams of sodium nitrate to sodium nitrite? How many grams of 
 lead would be required for the same purpose? At the current 
 
154 NITROGEN 
 
 market price which would be cheaper? Would the cost of the 
 manufacture depend primarily on the cost of the metal used? 
 
 5. How many pounds of ammonium nitrate will be required to 
 furnish 50 pounds of nitrous oxide? How many liters of the gas 
 will this give, counting 453 grams to the pound? 
 
 6. How many pounds of aqua ammonia of sp. gr. 0.90 and con- 
 taining 25 per cent of the gas and how many pounds of nitric 
 acid of sp. gr. 1.42 and containing 70 per cent of nitric acid will be 
 required to furnish the ammonium nitrate for the last problem? 
 The specific gravities are given as a matter of information and are 
 not to be used in the calculation. 
 
CHAPTER XIV 
 
 AIR; THE NOBLE GASES; GROUP ZERO 
 
 Composition of the Air. Oxygen. If air is left for some 
 time in contact with phosphorus, the oxygen will gradually 
 combine with the phosphorus leaving the ^^ 
 
 other gases contained in the air unchanged. 
 
 A quite accurate determination of the 
 per cent of oxygen in air may be made by ; ~~| 
 
 the apparatus shown in Fig. 30. A volume 
 of air is measured over water in the gradu- 
 ated tube (eudiometer), and then a piece 
 of phosphorus on the end of a wire is in- 
 serted. After some hours the phosphorus 
 is removed and the decrease in volume 
 compared with the original volume will give 
 the per cent of oxygen in the air. 
 
 Composition of the Air. Carbon Di- 
 oxide. If air is drawn through lime water 
 (a solution of calcium hydroxide, Ca(OH) 2 ), 
 the latter slowly becomes turbid from the 
 formation of calcium carbonate, CaC0 3 : 
 
 Ca(OH) 2 H- CO 2 = CaCO 3 + H 2 O 
 
 The same test has been applied to show 
 the presence of carbon dioxide in the gas 
 formed by burning charcoal in oxygen (p. 
 10). The presence of carbon dioxide in 
 the breath can be shown by blowing air from the lungs 
 through lime water, which will quickly grow turbid from 
 the separation of calcium carbonate. 
 
 155 
 
 FIG. 30. 
 
156 AIR; THE NOBLE GASES; GROUP ZERO 
 
 Since coal, wood, oil, gasolene and other compounds con- 
 taining carbon are constantly burning all over the world 
 and since the breath of millions of men and animals is con- 
 stantly exhaled into the air, it is evident that some carbon 
 dioxide must always be present in the atmosphere. 
 
 Some carbon dioxide escapes into the air from volcanoes. 
 Large quantities are formed from the decay of organic 
 matter in the soil and elsewhere under the influence of 
 bacteria. 
 
 If no means were provided for the removal of carbon 
 dioxide from the air it is evident that in the course of time 
 the amount in the air must very greatly increase. This 
 is prevented by the growth of trees "and plants, which secure 
 their supply of carbon from the carbon dioxide of the air 
 and constantly return the oxygen of the carbon dioxide to 
 the air. The energy for this process of reduction comes 
 from the sunlight. 
 
 The amount of carbon dioxide normally present in the air 
 is about 3 parts in 10,000 or 0.03 per cent by volume. 
 
 Composition of the Air. Water Vapor. The proportion 
 of oxygen and of carbon dioxide in air out of doors varies 
 very little indeed, but the per cent of water vapor varies 
 between wide limits. In a tropical climate air saturated 
 with water vapor might contain 5 per cent or more by 
 volume of the vapor, while in winter with the temperature 
 at 10 below zero the volume could not be as much as 0.3 
 per cent unless the air were supersaturated. Evidence 
 for the presence of water in the air is so common that no 
 experiments need be given to demonstrate it. 
 
 Composition of the Air. Argon. As long ago as 1785 
 the English chemist, Cavendish, mixed air with an excess 
 of oxygen, passed the sparks from a frictional electrical 
 machine through the mixture and absorbed the oxides of 
 nitrogen with an alkali. After continuing the experiment 
 for a long time he absorbed the oxygen which was left with 
 
DISCOVERY OF ARGON 157 
 
 " liver of sulfur" and found that the gas which remained was 
 not more than Jj[20 f ^ ne volume of the original air. For 
 more than 100 years this experiment was taken as proof 
 that the gas remaining when oxygen, carbon dioxide and 
 water vapor have been removed from the air is pure nitrogen. 
 
 In 1894 Lord Rayleigh, another Englishman, undertook 
 to make a very careful determination of the weight of a 
 liter of nitrogen. For this purpose he prepared and weighed 
 the gas which remains when oxygen, carbon dioxide and 
 water vapor are removed from common air. This gas was 
 then supposed to be pure nitrogen. For some reason Lord 
 Rayleigh thought that it would be of interest to prepare nitro- 
 gen from ammonia or from some other compound of nitrogen. 
 To his surprise he found that a liter of pure nitrogen pre- 
 pared by a chemical method weighs about 5 milligrams less 
 than a liter of the gas which he had prepared from air. This 
 was a much greater difference than could be accounted for 
 by the ordinary errors of the determination. Accordingly 
 Lord Rayleigh, with the assistance of Sir William Ramsay, 
 repeated the Cavendish experiment and obtained a gas 
 which is very much heavier than nitrogen and which they 
 called argon. 
 
 The Noble Gases. Argon proved to be not only a new 
 and hitherto unknown element, but an element wholly 
 different in its properties from all other elements then known. 
 The most distinctive property of the element is that it 
 seems to have no chemical affinity. It has not been found 
 possible to prepare a compound of argon with any other 
 element. 
 
 Within a few years Sir William Ramsay discovered five 
 other gaseous elements with similar properties. Gold 
 and platinum have been called noble metals because they 
 do not tarnish in the air and no single, ordinary acid will 
 attack them. Because of their lack of chemical affinity 
 the elements of this family have been called, for a similar 
 
158 AIR; THE NOBLE GASES; GROUP ZERO 
 
 reason, the noble gases. They are helium, neon, argon, 
 krypton, xenon and niton. 
 
 The relation of the elements of this group to the halogen 
 family on the one side and to the alkali metals on the other 
 is shown in the following table: 
 
 He 4 Li 7 
 
 F 19 Ne 20 Na 23 
 
 Cl 35.5 A 40 K 39 
 
 Br 80 Kr 83 Rb 85.8 
 
 I 127 Xe 130 Cs 133 
 
 Nt 222 
 
 Helium was discovered by Lockyer in the atmosphere of 
 the sun in 1868. It was not known that it exists on the 
 earth till it was discovered by Ramsay in some minerals, 
 a few years after the discovery of argon. Some years later 
 it was shown that helium is formed by the spontaneous 
 decomposition of radium and of other radioactive elements. 
 As our knowledge of the composite nature of the elements 
 increases it becomes more and more probable that helium 
 forms a part of the atoms of many different elements. 
 
 Neon, Krypton and Xenon are present in very small 
 amounts in the air. 
 
 Niton is formed by the decomposition of radium, helium 
 being formed at the same time. Niton in turn, decom- 
 poses very much more rapidly than radium and it is possi- 
 ble to obtain only very minute quantities of the gas. The 
 density was determined by Ramsay with the use of only a 
 few cubic millimeters. The element was called at first ra- 
 dium emanation. It is powerfully radioactive (p. 269). 
 
 Air is a Mixture. That the elements present in air are not 
 combined with each other is evident from the following facts: 
 
 1. Wherever chemical combination occurs heat is gen- 
 erated or absorbed. There is no change in temperature 
 when oxygen, nitrogen and argon are mixed. 
 
VENTILATION 159 
 
 2. If water is boiled, a gas resembling air escapes from it 
 but when the gas is examined it is found that it contains a 
 larger per cent of oxygen than air does. It has been shown 
 that for such gases as oxygen and nitrogen the weight of the 
 gas absorbed by a given volume of water varies directly 
 with the pressure (Henry's law). Oxygen is more soluble 
 than nitrogen, in water, and when we take account of the 
 solubility of the two gases and of their relative pressures 
 in air (0, 0.21; N, 0.78) the composition of the gas expelled 
 by boiling water agrees with that which is calculated on the 
 supposition that the oxygen and nitrogen are simply mixed 
 together. If they were combined, the composition of the 
 gas expelled from the water would be the same as the 
 composition of the air. 
 
 3 . The weight of a liter of dry air is almost exactly that cal- 
 culated on the supposition that the four gases, oxygen, nitro- 
 gen, carbon dioxide and argon, mix without combination. 
 
 Weight of Proportion in 
 
 one liter the air 
 
 Oxygen 1.429 X 0.2095 = 0.2994 
 
 Carbon dioxide .... 1 . 9768 X . 0003 = . 0006 
 
 Argon 1.7828 X 0.0094 = 0.0168 
 
 Nitrogen 1.2507 X 0.7808 = 0.9765 
 
 1.2933 
 
 The weight of a liter of air by direct determination is 
 1.2928 grams. 
 
 Ventilation. It was formerly supposed that the accumu- 
 lation of carbon dioxide from the breath is the chief source 
 of danger in poorly ventilated rooms. It has now been 
 demonstrated that this is not true and that the amount of 
 carbon dioxide present in the air of even badly ventilated 
 rooms is never great enough to cause any injury to human 
 beings. 
 
 In spite of this it must be considered as well established 
 that lack of ventilation in factories, offices and dwellings 
 
160 AIR; THE NOBLE GASES; GROUP ZERO 
 
 is a frequent cause of disease. It is also very well estab- 
 lished that abundance of fresh air, secured by life out of 
 doors both by night and day, combined with a nourishing 
 diet, furnishes the best hope of recovery from incipient 
 tuberculosis. 
 
 While exhaled carbon dioxide is not in itself harmful, it 
 furnishes the best means of determining whether a room is 
 properly ventilated or not. The amount of the gas should 
 not exceed 0.07 per cent by volume. To secure this stand- 
 ard of ventilation 55,000 liters or 2000 cubic feet of fresh 
 air will be required each hour for each person in the room. 
 
 SUMMARY 
 
 The per cent of oxygen in the air may be determined by 
 absorbing the oxygen with phosphorus. 
 
 The air receives carbon dioxide from the breath of men 
 and animals, from the combustion and decay of organic 
 matter and from volcanoes. 
 
 Carbon dioxide is removed from the air by growing plants. 
 The amount in ordinary air is 0.03 per cent or about 3 
 parts in 10,000. 
 
 Water vapor in the air varies between wide limits. 
 
 Argon was discovered by Lord Rayleigh in an attempt to 
 determine the exact weight of a liter of nitrogen. 
 
 The Zero Group of elements, sometimes called the noble 
 gases, consists of helium, neon, argon, krypton, xenon and 
 niton. 
 
 Helium was discovered in the sun by Lockyer. It was 
 discovered in some minerals later. 
 
 Niton is formed by the disintegration of radium. 
 
 Air is believed to be a mixture because no heat is evolved 
 when nitrogen and oxygen are mixed, because water dis- 
 solves the oxygen and nitrogen independently, and because 
 the weight of a liter corresponds to the calculated weight, 
 on the supposition that it is a mixture. 
 
EXERCISES. AIR 161 
 
 Good ventilation of rooms in which people live, and es- 
 pecially of sleeping rooms, office rooms and audience rooms 
 is very important. The carbon dioxide in such rooms 
 should not exceed 0.07 per cent, though the carbon dioxide 
 does not seem to be, in itself, harmful. 
 
 EXERCISES 
 
 1. What is the weight of a gram-molecular volume of air? 
 
 2. How many liters of air will be required to burn 32 grams of 
 sulfur? 31 grams of phosphorus? 12 grams of carbon? 56 
 grams of iron? 
 
 3. What facts of common experience demonstrate the presence 
 of water in the air? 
 
 4. What weight of water vapor do 22.4 liters of air at 20 con- 
 tain, assuming that the weight of a gram-molecular volume of 
 water vapor is 18 grams under normal conditions? What per 
 cent of water vapor does the air contain by weight? Is moist air 
 heavier or lighter than dry air? 
 
 5. A sample of bituminous coal has the following composition: 
 
 Carbon 66 . 25 per cent 
 
 Hydrogen (other than moisture) ... 4 . 25 per cent 
 
 Oxygen 8 . 00 per cent 
 
 Nitrogen 1 . 50 per cent 
 
 Moisture 9 . 00 per cent 
 
 Ash 11 .00 per cent 
 
 How many grams of oxygen will be required to burn one pound 
 (453 grams) of the coal? How many liters of oxygen? How 
 manv liters of air? 
 
CHAPTER XV 
 THE PERIODIC SYSTEM 
 
 When the elements are arranged in the order of their 
 atomic weights it is found that elements of similar properties 
 occur at regular intervals. This fact has already been used 
 as the basis for the classification of the non-metallic elements 
 which have been studied in the preceding pages. A classi- 
 fication of all of the elements on the basis of their atomic 
 weights is given in the accompanying tables: 
 
 For an elementary study of these tables the following 
 relations are of most importance: 
 
 1. Groups and Periods. The elements in a perpendicular 
 column are classed together as a group. In the first table 
 there are nine groups numbered to VIII but it. should 
 be noticed that Group VIII includes three sets of elements, 
 of three in each set, and that it differs in this respect from 
 the other groups, which contain only a single element 
 at a given point, in the earlier part of the table. 
 
 The elements in a horizontal line are classed together as 
 periods, but it is necessary, here, to distinguish between 
 the two "short periods" which include the first two lines 
 of the first table and two "longer periods" which include 
 the next four lines of the first table, or the third and fourth 
 lines of the second table. The second table is designed 
 to bring out these long periods more clearly and to show 
 that in the first table after chlorine the alternate elements 
 of each group resemble each other more closely than the 
 successive elements in the group. Thus potassium, rubid- 
 ium and caesium, of Group I, are closely alike and are usually 
 
 162 
 
COMPOSITE NATURE OF ATOMS 163 
 
 spoken of as the first division of the group, while copper, 
 silver and gold are alike and form the second division of 
 the group. 
 
 In the rest of the table the fifth and sixth periods are still 
 longer than the third and fourth as shown in the second 
 table. If the rare earth elements are included, as they 
 certainly must be if we are not to lay ourselves open to the 
 objection that facts have been distorted to fit the theory of 
 the table, the fifth period must include all elements from 
 xenon to bismuth, and only four elements, whose atomic 
 weights have been determined, are known for the sixth 
 period. 
 
 2. Composite Nature of Atoms. Long before the radio- 
 active elements (p. 268) were discovered many chemists 
 considered that the periodic system points clearly toward 
 a belief in the composite nature of the atoms of the elements. 
 The discovery of the disintegration of the atoms of radio- 
 active elements has given almost conclusive proof of the 
 truth of this view and has shown that atoms of helium and 
 electrons (atoms of negative electricity) are constituent 
 parts of the atoms of some of these elements. For a long 
 time some chemists have proposed the hypothesis that 
 atoms of hydrogen are also constituent parts of the atoms 
 of the other elements, and some facts point rather strongly 
 toward such a view. If we suppose, as seems reasonable, 
 that all atoms are composite, in order to account for the 
 atomic weights it seems necessary to suppose that particles 
 which are intermediate in weight between helium (4) and 
 electrons (1/1800 the weight of hydrogen atoms) must 
 enter into the composition. 
 
 Thus far, however, no one has shown that hydrogen 
 atoms are found among the disintegration products of any 
 element. 
 
 3. The position of an element in the system indicates at 
 least one of its valences, especially its highest valence toward 
 
164 
 
 THE PERIODIC SYSTEM 
 
 
 
 
 
 C 
 
 s? 
 
 fctf 
 
 lo. 
 oS 
 
 II 
 
 10 
 
 
 
 PQ ^H 
 
 o 
 
 tf 
 
 u 
 
 - 
 
COMPOSITE NATURE OF ATOMS 
 
 165 
 
 i, SS 
 
 ff 
 
 S 
 
 58 OT 
 
 o g 
 
 ss 
 
 s 
 
 Ml 
 
 c, a 
 
 a * 
 
 H 
 
 rt o 
 o 
 
 
166 THE PERIODIC SYSTEM 
 
 oxygen (positive valence) and, for the non-metallic ele- 
 ments, the valence toward hydrogen (negative valence). 
 This will be clear from the following formulas of compounds 
 of the elements of the second period. 
 
 Group I Group II Group III Group IV Group V Group VI Group VII 
 
 Na Mg=O O O O O O O 
 
 \ /" ./ /" / II / 
 
 Cl=( 
 
 O Al C N=O S=O C1=O 
 
 / \ \ \ \ \ 
 
 Na O O O O O 
 
 Al N=O C1=O 
 
 \ \ II \ 
 
 O O O O 
 
 H H H H 
 
 N H S Cl H 
 
 i-H 
 
 H H H 
 
 It should be noticed that the sum of the valences toward 
 oxygen and toward hydrogen is eight. This is doubtless 
 connected in some way with the structure of the atoms, but 
 our knowledge of that structure is still too indefinite to give 
 any indication of the basis for this or, indeed, for many 
 other puzzling facts. 
 
 4. In the periods metallic properties decrease and non- 
 metallic properties increase from left to right, that is, 
 with increasing atomic weight. In the groups, on the 
 other hand, metallic properties increase and non-metallic 
 properties decrease from top to bottom. The result is 
 that the non-metallic elements are found in the upper right- 
 hand corner of the table, above the dotted line. Strictly 
 speaking, the Zero Group has no chemical properties and 
 we cannot classify the elements of that group as either 
 metallic or non-metallic. It would, perhaps, be more 
 logical to place the group to the right of the table as alter- 
 nate with the elements of Group VIII. With such an 
 
g 3 
 
168 THE PERIODIC SYSTEM 
 
 arrangement there would be two divisions in Group VIII, 
 as in the other groups. 
 
 5. The physical properties of the elements, such as 
 melting points and specific gravity, are closely connected 
 with their positions in the table. The relation between 
 the melting points and the positions in the table will be 
 seen by an examination of the melting points given in the 
 second table. The use of tungsten (W) for electric lights 
 was suggested by the position of the element in the tab'le. 
 
 The relation of the specific gravity of the elements to 
 their positions in the table is shown in Fig. 31. The 
 atomic volume of an element is the volume occupied by one 
 gram atom. Thus the specific gravity of potassium (K) 
 is 0.862 and a gram atom of the element weighs 39.1 grams. 
 
 39 1 
 The volume occupied by a gram atom will be Q g62 = 
 
 45.4 cc. and this is the atomic volume of potassium. 
 
 6. Atomic Numbers. An examination of the periodic 
 table reveals the fact that three elements, argon (A), 
 cobalt (Co) and tellurium (Te), are not placed in the table 
 in accordance with their atomic weights. As the properties 
 of these elements leave no question but that they are prop- 
 erly placed, these exceptions have caused some misgiving 
 as to the validity of the table. This has led to a number of 
 very careful determinations of the atomic weights of these 
 elements and especially of tellurium. These determinations 
 have confirmed the values for the atomic weights which are 
 given and facts of this sort are always accepted by chemists 
 no matter how much they may conflict with theories or 
 systems which are in vogue. 
 
 In comparatively recent times it has been discovered by 
 means of X-ray spectra that each element has a charac- 
 teristic property, called its atomic number. This atomic 
 number is approximately one-half of the atomic weight and 
 appears to be connected with the structure of the atom in 
 
EXERCISES 169 
 
 such a manner that it is a more fundamental charac- 
 istic than the atomic weight. The atomic numbers of 
 argon, cobalt and tellurium agree with the positions assigned 
 them in the periodic table. 
 
 SUMMARY 
 
 Elements are classified in a table, called the periodic 
 system, which is based on bheir atomic weights. 
 
 There are nine groups of elements and each group except 
 the Zero Group and Group VIII is separated into two divi- 
 sions. 
 
 There are two " short periods," two "long periods," 
 a fifth very long period and a sixth period for which only a 
 very few elements are known. 
 
 The valences, metallic or non-metallic properties, melting 
 points and specific gravities of the elements are 1 closely 
 connected with the positions of the elements in the periodic 
 system. 
 
 The atomic volume of an element is the volume in cubic 
 centimeters occupied by one gram atom of the element. 
 
 The atomic numbers of the elements," determined by 
 means of their X-ray spectra, are even more characteristic 
 than their atomic weights. 
 
 EXERCISES 
 
 1. Prepare tables of the oxides, acids, chlorides and hydrogen 
 compounds of the elements which have been studied, showing 
 the relationships between them. 
 
 2. The specific gravity of copper is 8.93. What is its atomic 
 volume? 
 
 3. From the positions in the periodic system what oxides are to 
 be expected from the following elements: Strontium, Sr, tita- 
 nium, Ti, vanadium, V, chromium, Cr, molybdenum, Mo, thor- 
 ium, Th, tantalum, Ta? 
 
 4. What elements besides carbon and tungsten have been used 
 for filaments of electric lights? 
 
CHAPTER XVI 
 
 GROUP V: PHOSPHORUS, ARSENIC, ANTIMONY AND 
 
 BISMUTH 
 
 The Nitrogen Family of Elements. Fluorine differs 
 very decidedly in many of its properties from the other- 
 members of the halogen family, chlorine, bromine and iodine, 
 and oxygen is quite different from sulfur, selenium and 
 tellurium. In a similar manner nitrogen differs from phos- 
 phorus, arsenic, antimony and bismuth. In spite of these 
 differences, however, there are enough resemblances so 
 that a comparison of the compounds and properties is of 
 much service in studying the elements of the group. 
 
 Occurrence of Phosphorus. Phosphorus is not found 
 free in nature. While free nitrogen is induced to combine 
 with other elements with difficulty and only under very 
 special conditions, ordinary phosphorus combines with 
 oxygen so readily and holds the element so firmly that it is 
 impossible that it should be liberated or exist free under any 
 ordinary conditions found in nature. It is found almost or 
 quite exclusively in the form of salts of phosphoric acid, 
 H 3 PO4. The most common salt is tricalcium phosphate, 
 Ca 3 (PO 4 )2. This salt is the principal mineral constituent 
 of bones and so of bone-ash, obtained by burning bones. It 
 is also found as a mineral, more or less pure, in North and 
 South Carolina, Tennessee, Georgia, Florida and in some 
 other states. Since phosphates are essential for the growth 
 of plants and some soils are deficient in them, large quan- 
 tities of these phosphates are used for fertilizing purposes. 
 Apatite is another mineral containing phosphorus. It is 
 a combination of calcium phosphate with the fluoride, 
 
 170 
 
PREPARATION AND PROPERTIES OF PHOSPHORUS 171 
 
 CaF 2 , or chloride, CaCl 2 , and has the formula Ca 5 (PO 4 ) 3 F 
 or Ca 5 (PO 4 )3Cl. The relation between the valence of 
 calcium and those of the phosphate radical, P0 4 , and of 
 fluorine should be noticed in this formula. 
 
 Preparation of Phosphorus. Phosphorus has been used 
 in the manufacture of matches for less than a century 
 but during that time the manufacture of the element has be- 
 come very important. For a long time it was obtained by 
 a complicated, tedious process in which calcium phosphate 
 was first treated with sulfuric acid and the phosphoric 
 acid or acid calcium phosphate formed was reduced by 
 heating it with carbon. During comparatively recent times 
 the high temperature which can be secured in an electric 
 furnace has made it possible to manufacture phosphorus by 
 heating a mixture of calcium phosphate, Ca 3 (PO 4 ) 2 , silica, 
 or sand, SiO 2 , and carbon, in the form of charcoal or coke: 
 
 2Ca 3 (PO 4 ) 2 + 6SiO 2 + IOC = 6CaSiO 3 + P 4 + 10CO 
 
 The silica, SiO 2 , combines with the calcium and oxygen 
 (lime), CaO, to form calcium silicate, CaSiO 3 , while the 
 carbon takes oxygen from the phosphorus. In writing the 
 equation notice that the formula of phosphorus is P 4 and 
 that two molecules of calcium phosphate are required to 
 give this. The rest follows logically. 
 
 Properties of Phosphorus. Ordinary phosphorus is a 
 wax-like, almost colorless solid when first prepared, but it 
 soon acquires a yellow or red color on exposure to the light. 
 Its kindling temperature is very low and it is kept under 
 water and cast into sticks. It boils at 290 and may be 
 easily distilled. Ordinary phosphorus is very poisonous. 
 
 Red Phosphorus. If heated at 240-250 for some time 
 in a closed ' vessel ordinary phosphorus is changed partly, 
 but not completely, to the allotropic form known as red 
 phosphorus. At a higher temperature red phosphorus is 
 converted into a vapor which condenses, on cooling, to the 
 
172 PHOSPHORUS, ARSENIC, ETC. 
 
 ordinary form. Red phosphorus does not take fire with 
 ordinary friction and it is not poisonous. It is used in the 
 manufacture of safety matches and for many purposes in 
 chemical laboratories. 
 
 Matches. Phosphorus was discovered nearly 250 years 
 ago but it was not used for the manufacture of matches till 
 1827. Persons still living can remember being sent to a 
 neighbor's to borrow coals of fire because of the difficulty 
 of starting a fire with flint, steel and tinder. During the 
 nineteenth century ordinary phosphorus was chiefly used 
 for making matches, though in European countries the or- 
 dinary phosphorus match came to be displaced by safety 
 matches, which must be ignited on a prepared surface and 
 offer less danger of accidental fires. The use of ordinary 
 matches in America is one of the reasons for the very much 
 greater losses by fire here than abroad. 
 
 Owing to the mortality due to phosphorus poisoning in 
 match factories, Congress in 1912 passed a law taxing 
 the use of ordinary phosphorus, and tetraphosphorus 
 trisulfide, P^Ss, is now used in place of it. This gives a 
 match which can be ignited by friction and still its manu- 
 facture does not endanger the health and lives of the work- 
 men. The match sticks are made of some wood which 
 does not readily retain a live coal after the flame is ex- 
 tinguished and this non-inflammable quality is further 
 increased by treatment with a dilute solution of magnesium 
 sulfate or alum. 
 
 Phosphine. When phosphorus is warmed with a con- 
 centrated solution of sodium hydroxide (Fig. 32), NaOH, 
 phosphine, PH 3 , a compound corresponding to ammonia, 
 NH 3 , is formed. The details of the reaction need not be 
 given here. A small amount of a second compound of 
 phosphorus and hydrogen, P2H 4 , is also formed and the 
 presence of this causes the phosphine prepared by the 
 method described to take fire on coming to the air. 
 
PHOSPHORUS PENTOXIDE 
 
 173 
 
 Phosphine combines with hydriodic acid, HI, to form the 
 salt phosphonium iodide, PH 4 I, which corresponds to the 
 similar salts ammonium iodide, NH 4 I, and ammonium 
 chloride, NH 4 C1, formed from ammonia. Phosphonium 
 iodide is far less stable than ammonium iodide, and the 
 compound is chiefly of interest because of the analogy 
 which it shows between phosphine and ammonia. 
 
 FIG. 32. 
 
 Phosphorus Pentoxide, or Phosphoric Anhydride. When 
 phosphorus is burned in dry air or in oxygen the oxide 
 formed has the composition represented by the formula 
 P 2 O 5 . The oxide may be converted into a vapor at a rather 
 high temperature and it has been found that the gram- 
 molecular volume of the vapor weighs about 284 grams, 
 from which it follows that the true formula is P 4 Oio and 
 not P2O 5 . The name phosphorus pentoxide is still used, 
 however, and the formula P 2 O 5 correctly represents the 
 composition of the compound, of course. The name phos- 
 phoric anhydride is strictly accurate and is, perhaps, better. 
 It is usually more convenient to use the formula P 2 Os in 
 
174 PHOSPHORUS, ARSENIC, ETC. 
 
 equations, just as equations are sometimes written in which 
 oxygen, hydrogen and other elements are involved without 
 taking account of the true formulas, 2 , H 2 , etc. 
 
 Phosphorus pentoxide has a very strong affinity for water 
 and it is the most perfect drying agent we have for the 
 gases which do not combine with it. 
 
 Phosphoric Acids. Phosphorus pentoxide may combine 
 with one, two or three molecules of water to form three dif- 
 ferent acids, which are all called phosphoric acids because 
 the phosphorus is in the same state of oxidation in each. 
 The addition or removal of water from a compound is 
 considered neither as an oxidation nor a reduction. The 
 relation of the three acids to phosphorus pentoxide is as 
 follows : 
 
 H 2 OP 2 O5 = 2HPO 3 , Metaphosphoric acid 
 2H 2 OP 2 O 5 = H 4 P 2 7 , Pyrophosphoric acid 
 3H 2 OP 2 5 = 2H 3 P0 4 , Orthophosphoric acid 
 
 The last of these acids, Orthophosphoric acid, is much 
 more common than either of the others and is always under- 
 stood when the name phosphoric acid, without a prefix, 
 is used. 
 
 Salts of Orthophosphoric Acid. As a tribasic acid 
 Orthophosphoric acid forms three classes of salts: acid 
 salts (p. 76), in which one or two of the hydrogen atoms 
 of the acid have been replaced, and normal salts, in which 
 all three atoms have been replaced by a metal. The salts 
 are distinguished as primary, secondary and tertiary, or, 
 more often, by the prefixes mono-, di- and tri-. The 
 latter refer to the numbers of hydrogen atoms replaced 
 from one molecule of the acid. Thus Ca(H 2 PO 4 )2 or CaH 4 - 
 (PO 4 ) 2 is monocalcium phosphate. CaHP0 4 is dicalcium 
 phosphate, although it has only a single atom of calcium 
 in the molecule. 
 
HYDROLYSIS OF SALTS OF PHOSPHORIC ACID 175 
 The relations will be clear from the following table: 
 
 Mpnosodium phosphate, NaH2PO4 Monocalcium phosphate, 
 Disodium phosphate, Na2HPO4 Dicalcium phosphate, CaHPO4 
 
 Trisodium phosphate, NasPO4 Tricalcium phosphate, Cas(PO4)z 
 
 Salts of pyrophosphoric acid and of metaphosphoric acid 
 are formed by heating the acid salts or mixed ammonium 
 salts derived from orthophosphoric acid. Thus either 
 monosodium phosphate, NaH 2 PO4, or sodium ammonium 
 phosphate, NaNH 4 HP0 4 , gives sodium metaphosphate, 
 NaPO 3 , when heated. The sodium ammonium phosphate 
 loses both ammonia and water. Give the formulas and 
 names of the salts which will be formed by heating the 
 following: CaHP0 4 , MgNH 4 PO 4 , KNH 4 HPO 4 ,CaH 4 (P0 4 ) 2 
 
 Hydrolysis of Salts of Phosphoric Acid. Phosphoric 
 acid, H 3 PO 4 , is a comparatively strong acid. In other 
 words, it is largely ionized in dilute solutions, giving hydro- 
 gen ions, H+, and monophosphate ions, H 2 PO 4 ~. To a 
 very slight extent it is further ionized to two hydrogen ions, 
 H + , H + , and the diphosphate ions, HPO 4 = . The diphos- 
 phate ions, HPO 4 = , also ionize to a very trifling extent into 
 hydrogen ions and triphosphate ions, PO 4 ~, but this last 
 ionization is probably less in amount than even the ioniza- 
 tion of water to hydrogen, H + , and hydroxide, OH~, ions. 
 
 If a solution of sodium hydroxide, NaOH, is added to a 
 dilute solution of phosphoric acid, H 3 PO 4 , in the proportion 
 of one molecule of the hydroxide for one molecule of the 
 acid, the hydrogen ions for the first ionization of the acid 
 will be removed by combination with the hydroxide ions of 
 the base. The solution will then contain sodium ions, 
 Na + , 'monophosphate ions, H 2 PO 4 ~, and a few hydrogen ions 
 H + , and diphosphate ions, HPO4 = . The solution will still 
 be faintly acid, but if more sodium hydroxide is added a 
 point is soon reached where the tendency of the diphos- 
 phate ions, HPO 4 = , to combine with hydrogen ions to form 
 monophosphate ions, H 2 PO 4 ~, will just balance the tend- 
 
176 PHOSPHORUS, ARSENIC, ETC. 
 
 ency of water to separate into hydrogen and hydroxide 
 ions and the numbers of hydrogen and of hydroxide ions 
 in the solution will be equal; the solution will then be 
 neutral. This point is reached long before the second 
 hydrogen atom of the phosphoric acid has been neutral- 
 ized. Any further addition of sodium hydroxide must 
 give an excess of hydroxide ions and the solution will have 
 an alkaline reaction. 
 
 If the compound disodium phosphate, Na 2 HPO 4 , is 
 dissolved in water it separates into sodium ions, Na + , Na+, 
 and diphosphate ions, HPO4 = . The diphosphate ions have 
 so strong a tendency to combine with hydrogen ions, how- 
 ever, that when there are many of them in a solution they 
 will combine with the hydrogen ions of the water, leaving 
 an excess of hydroxide ions. This will give the solution an 
 alkaline reaction-. 
 
 A decomposition of a salt or compound by the ions of 
 water, in this way, is called hydrolysis (from i55wp, water, 
 and Xuco, I loose) because it is caused by the loosening 
 effect of the water. 
 
 Other Acids of Phosphorus. Two other acids of phos- 
 phorus, phosphorous acid, H 3 PO 3 , and hypophosphorous 
 acid, H 3 PO 2 , are fairly well known. Phosphorous acid 
 is bibasic. No salts have been prepared in which all three 
 of its hydrogen atoms have been replaced. Hypophos- 
 phorous acid is monobasic. What will be the formulas of 
 the normal sodium and calcium salts of these acids? 
 
 Chlorides of Phosphorus. Hydrolysis. When phos- 
 phorus burns in chlorine either phosphorus trichloride, 
 PG1 3 , or phosphorus pentachloride, PC1 5 , may be formed 
 according to the conditions of the experiment. These 
 chlorides are decomposed (hydrolyzed) by water with the 
 formation of hydrochloric acid and either phosphorous 
 acid or phosphoric acid: 
 
ARSENIC 177 
 
 Cl H - OH HO 
 P Cl + H - OH = HO P + 3HC1 
 
 Cl H - OH HO 7 
 
 Cl H-O-H H-Ov 
 /Cl H-0-H=H- 0-^p = o + 5HC1 
 P/.C1 H-O-H H-O/ 
 \SC1 H-O 
 
 H/ 
 
 or PC1 3 -f 3H 2 O = H 3 PO 3 + 3HC1 
 
 and PC1 5 + 4H 2 O = H 3 PO 4 + 5HC1 
 
 Arsenic, As, 74.96. 'Arsenic is present in considerable 
 amounts in some of the copper ores in the West and in 
 smelting these ores the arsenic is oxidized to arsenic trioxide, 
 which is produced in enormous quantities at the same 
 time. The deposit of the arsenic trioxide on vegetation in 
 the neighborhood of the smelters has caused considerable 
 trouble from the poisoning of cattle. Arsenic is also found 
 in arsenical pyrites, a mineral called mispickel and having the 
 formula FeAsS. Arsenic may be obtained either by heating 
 mispickel or by reducing the trioxide with carbon. When 
 not tarnished arsenic is a steel-gray, brittle metal which 
 may be easily powdered. When heated it volatilizes with- 
 out melting. In the air the vapors burn to the trioxide, 
 As 2 O 3 . 
 
 Arsenic trioxide is a white, crystalline powder, formed by 
 burning the metal. It is nearly or quite tasteless and is a 
 violent poison. It is used as a ratsbane and has often been 
 used in cases of criminal poisoning. It is also used as a 
 medicine. The methods of chemical analysis make it 
 possible to detect a very minute quantity of the element. 
 
 Arsine. If a soluble compound of arsenic is introduced 
 into a generator in which hydrogen is prepared from zinc 
 and hydrochloric or sulfuric acid, the arsenic is reduced to 
 arsine, AsH 3 . The arsine escapes with the hydrogen. It 
 
 12 
 
178 
 
 PHOSPHORUS, ARSENIC, ETC. 
 
 imparts to the hydrogen flame a pale blue color and metallic 
 arsenic is deposited on a piece of cold porcelain held in 
 the flame, very much as carbon is deposited from a candle 
 flame. If the hydrogen containing arsine is passed through 
 a heated tube, the arsine is decomposed and a mirror of 
 arsenic is deposited on the walls of the tube. By the arrange- 
 ment shown in Fig. 33 it is possible to detect 1/1000 of a 
 milligram of arsenic. It is, of course, necessary to use zinc 
 and acid, which are entirely free from arsenic. 
 
 FIG. 33. 
 
 Arsenic trichloride is a colorless liquid which may be pre- 
 pared by the direct union of arsenic and chlorine. It is 
 hydrolyzed by water in the same way as phosphorus tri- 
 chloride and other chlorides of non-metallic elements. 
 Arsenic trioxide separates from the solution. Some arse- 
 nic remains in the solution, however, and the conduct of 
 the solution toward hydrogen sulfide indicates that a part 
 of this is still in the form of the chloride. In the reactions 
 represented by the equations: 
 
 2AsCl 8 + 6HOH <=* 6HC1 + 2As(OH) 3 <= As 2 O 3 + 3H 2 O 
 
 the equilibrium is far to the right and the course of the 
 reaction in that direction is also favored by the slight solu- 
 bility of the arsenic trioxide. 
 
 Arsenious Acid. Just as nitrous acid, HNO 2 , easily 
 decomposes into nitrous anhydride, N 2 Os, and water, arse- 
 nious acid, HsAsOa, decomposes into arsenic trioxide, 
 
ARSENIC ACID 179 
 
 As 2 03, the anhydride of the acid, and water. Salts of the 
 acid are known, however, such as silver arsenite, Ag 3 AsO 3 , 
 and Paris green, Cu(C 2 H 3 O2)2 . Cu 3 (As0 3 ) 2 . The latter is 
 a double salt of copper acetate and copper arsenite. It is 
 used as a green pigment for painting and as a poison for 
 potato bugs and other insects. 
 
 Arsenic Acid. When arsenious oxide is warmed with 
 nitric acid it is oxidized to arsenic acid, H 3 As0 4 , while the 
 nitric acid is reduced to nitrous anhydride, N2O 3 . This is, 
 indeed, the most -convenient method of preparing nitrous 
 anhydride : 
 
 As 2 O 3 + 2HNO 3 + 2H 2 O = 2H 3 AsO 4 + N 2 O 3 
 
 Arsenic acid is tribasic. The two most interesting salts 
 are trisilver arsemate, Ag 3 AsO 4 , and magnesium ammonium 
 arsenate, MgNH 4 As0 4 . What will be the formula and 
 what is the name of the compound formed by heating the 
 latter (see p. 175)? 
 
 Arsenic trisulfide, As 2 S 3 , is formed by passing hydrogen 
 sulfide into a solution of arsenic trioxide, As2O 3 , in hydro- 
 chloric acid. The formation of the compound indicates 
 the presence of arsenic trichloride, AsCl 3 , in the solution. 
 Arsenic trisulfide is found as a mineral called orpiment, and 
 is sometimes used as a yellow pigment by artists. 
 
 Arsenic trisulfide dissolves easily in a solution of am- 
 monium sulfide (NH 4 )2S, giving ammonium sulf arsenite, 
 (NH 4 ) 3 AsS 3 : 
 
 As 2 S 3 + 3(NH 4 ) 2 S = 2(NH 4 ) 3 AsS 3 
 
 Ammonium sulf arsenite receives its name from the simi- 
 larity to ammonium arsenite (NH 4 ) 3 AsO 3 . It is used in 
 qualitative analysis to separate arsenic from copper, lead and 
 other metals whose sulfides do not form similar compounds. 
 
 Antimony, Sb, 120.2. Antimony occurs chiefly as the 
 mineral stibnite, which is antimony trisulfide, Sb 2 S 3 . 
 
180 PHOSPHORUS, ARSENIC, ETC. 
 
 Precipitated antimony trisulfide is orange colored but the 
 mineral stibnite is black. Metallic antimony is obtained 
 from stibnite by roasting it in the air to convert the metal 
 into the oxide and then reducing the oxide with carbon: 
 
 2Sb 2 S 3 + 9O 2 = 2Sb 2 O 3 + 6SO 2 
 2Sb 2 O 3 + 3C = 4Sb + 3C0 2 
 
 Antimony is a brittle metal which may be easily pulver- 
 ized. When heated in the flame of a blowpipe on charcoal 
 it melts easily and burns slowly to antimony trioxide, 
 Sb 2 O 3 , which escapes as a white cloud. 
 
 In its metallic luster and conductivity for electricity 
 antimony is clearly metallic. In its brittleness it resembles 
 the non-metals. It is sometimes called a half-metal. 
 
 An alloy containing lead and antimony is used for type 
 metal. The antimony gives hardness and also causes the 
 metal to expand in solidifying, giving sharp, clear-cut type. 
 An alloy with lead, copper and a little bismuth is made as an 
 antifriction metal for bearings in machinery. One of these 
 alloys is called Babbitt metal. 
 
 Stibine, SbH 3 , may be prepared in the same manner as 
 arsine, AsH 3 , and resembles it closely in its properties. It 
 decomposes at a lower temperature than arsine and gives 
 a blacker, more sooty spot on porcelain. It is to be noticed 
 that neither arsine or stibine combines with acids, differing 
 in this respect from ammonia and phosphine. 
 
 Antimony trichloride, SbCl 3 , may be prepared by dis- 
 solving either the trioxide, Sb 2 O 3 , or the trisulfide, Sb 2 S 3 , 
 in concentrated hydrochloric acid. The solution may be 
 evaporated and the trichloride distilled. This is in very 
 marked contrast with the conduct of arsenic trichloride and 
 illustrates the increase in metallic and decrease in non- 
 metallic properties with increasing atomic weight in the 
 group. Chlorides of non-metals are hydrolyzed by water. 
 Chlorides of metals dissolve and ionize in water. 
 
BISMUTH 181 
 
 Antimony trichloride, however, shows also the properties 
 of a non-metallic chloride. If the solution is diluted it is 
 hydrolyzed in part and antimony oxychloride, SbOCl, is 
 precipitated : 
 
 SbCl 3 + H 2 O = SbOCl + 2HC1 
 
 Oxides of Antimony. Antimony forms three oxides, 
 Sb 2 O 3 , Sb 2 O 4 and Sb 2 O 5 , which correspond to the three 
 highest oxides of nitrogen. 
 
 Antimony Hydroxide, Sb(OH) 3 or Antimonious acid, 
 H 3 SbO 3 , may act either as a base or an acid: 
 
 Sb(OH) 3 + 3HC1 = SbCl 3 + 3H 2 
 H 3 SbO 3 + NaOH = NaH 2 SbO 3 + H 2 O 
 
 Sodium antimonite 
 
 Here, again, it is seen that the conduct of the compounds 
 of antimony places it midway between the metals and the 
 non-metals and that it is partly both a metal and a non- 
 metal. 
 
 Antimony trisulfide is formed as an orange-colored pre- 
 cipitate when hydrogen sulfide is passed into an acid solution 
 of an antimony salt. It dissolves in ammonium sulfide, 
 forming ammonium sulfantimonite, (NH 4 ) 3 SbS 3 , corre- 
 sponding to arsenic sulfarsenite, (NH 4 ) 3 AsS 3 . 
 
 Bismuth, Bi, 208. Bismuth is found mostly in the free 
 state in nature. It is a brittle metal and has a slightly 
 reddish cast. It is used in a variety of alloys, chiefly to 
 lower the melting point. Thus it is used in the stereotype 
 metal, employed in printing daily papers, to lower the 
 melting point so far that it does not injure the papier- 
 mache mold. Wood's metal (Bi 4 parts, Pb 2 parts, Sn 
 1 part, Cd 1 part) melts at 60.5. A teaspoon made from 
 the metal will melt in a hot cup of tea. Alloys similar 
 to Wood's metal are used for safety plugs in steam boilers 
 and in automatic sprinklers for protection against fire and 
 for the safety-fuses in electrical work. A small amount 
 
182 PHOSPHORUS, ARSENIC, ETC. 
 
 of bismuth has been found to improve the quality of Babbitt 
 metal. 
 
 Compounds of Bismuth. Bismuth forms no compound 
 corresponding to arsine, AsH 3 , and stibine, SbH 3 . Bismuth 
 trioxide, Bi 2 O3, does not dissolve in alkalies, and bismuth 
 sulfide, Bi 2 S 3 , does not dissolve in ammonium sulfide. 
 In these regards bismuth is more distinctly metallic than 
 any other member of the group. The chloride, BiCl 3 , 
 is hydrolyzed by water to the oxychloride, BiOCl, and 
 the nitrate, Bi(NO) 3 , to the basic nitrate, Bi(OH) 2 N0 3 , or 
 to bismuthyl nitrate, BiON0 3 : 
 
 /OH 
 
 Bi(NO 3 ) 3 + 2HOH = Bi OH + 2HNO 3 
 
 /OH Q 
 
 Bi OH = Bi/ + H 2 
 \N0 3 XN 3 
 
 This basic nitrate, which varies in composition according 
 to the method of preparation, is used in medicine and as a 
 face powder. Pharmacists still often use for it the anti- 
 quated name "subnitrate of bismuth." 
 
 SUMMARY 
 
 Nitrogen is quite different from the other elements of 
 Group V, phosphorus, arsenic, antimony and bismuth. 
 
 Phosphorus is found as tricalcium phosphate in bone- 
 ash and as a mineral; also as apatite. 
 
 Phosphorus is prepared by heating calcium phosphate, 
 silica and carbon in an electric furnace. 
 
 Phosphorus exists as white and red phosphorus. The 
 former is very poisonous. 
 
 Tetraphosphorus trisulfide is used in making matches. 
 Red phosphorus is used for safety matches. 
 
SUMMARY. PHOSPHORUS, ETC. 183 
 
 Phosphine is prepared by warming phosphorus with a 
 solution of sodium hydroxide. So prepared it .takes fire 
 in the air. It combines with hydriodic acid. 
 
 Phosphorus pentoxide is prepared by burning phosphorus 
 in dry air. It is an effective drying agent. 
 
 There are three phosphoric acids, orthophosphoric, 
 pyrophosphoric and metaphosphoric. 
 
 Orthophosphoric acid is a tribasic acid but only the first 
 hydrogen atom ionizes to any considerable extent. 
 
 Bimetallic and trimetallic phosphates are hydrolyzed 
 by water because the second and third hydrogen atoms 
 have only faint acid properties. Solutions of these salts 
 have an alkaline reaction. 
 
 Hydrolysis is the decomposition of a substance by water, 
 especially by the hydrogen and hydroxide ions of the water. 
 
 Phosphorous acid is dibasic and hypophosphorous acid 
 is monobasic. 
 
 Phosphorus trichloride and phosphorus pentachloricle are 
 hydrolyzed to phosphorous and phosphoric acids and 
 hydrochloric acid. 
 
 Arsenic is found in copper ores and as mispickel, arsenical 
 pyrites. 
 
 Arsenic trioxide is used as a ratsbane. 
 
 Arsine is formed when an arsenic compound is introduced 
 into a hydrogen generator. It is used in the detection of 
 arsenic. 
 
 Arsenic trichloride is hydrolyzed to arsenious oxide and 
 hydrochloric acid. 
 
 Arsenious acid is known only in very dilute solutions. 
 
 Paris green is a double salt of arsenious acid and acetic 
 acid. 
 
 Arsenic acid is tribasic. The silver salt and the magne- 
 sium ammonium salts are of interest. 
 
 Arsenic trisulfide is lemon yellow. It dissolves in am-^ 
 monium sulfide to give ammonium sulfarsenite. 
 
184 PHOSPHORUS, ARSENIC, ETC. 
 
 Antimony occurs as stibnite. It is prepared by roasting 
 the ore and reducing the oxide formed with carbon. 
 
 Antimony is brittle but more metallic than arsenic. It 
 is used in type metal and Babbitt metal. 
 
 Stibine is prepared in the same manner as arsine. It 
 decomposes at a lower temperature. 
 
 Antimony trichloride is hydrolyzed to antimony oxy- 
 chloride. 
 
 Antimony forms three oxides, the trioxide, tetroxide 
 and pentoxide. 
 
 Antimony hydroxide has both acid and basic properties. 
 Antimony trisulfide is orange. It dissolves in ammonium 
 sulfide to give antimony sulfantimonite. 
 
 Bismuth is found free. It is used in stereotype metal, 
 fusible metal and in anti-friction alloys. 
 
 Bismuth trioxide does not dissolve in alkalies and the tri- 
 sulfide does not dissolve in ammonium sulfide, indicating 
 that bismuth is the most metallic element of the group. 
 
 Bismuth chloride is hydrolyzed to the oxychloride and 
 bismuth nitrate to a basic nitrate. The latter is used in 
 face powders. 
 
 EXERCISES 
 
 1. Prepare a table of the hydrogen compounds, oxides, chlorides, 
 acids, and salts of the sulfur acids, of the elements of Group V. 
 
 2. Are ammonia, phosphine, arsine and stibine heavier or lighter 
 than air and in approximately what proportions? 
 
 3. In what proportion should phosphine and oxygen be mixed 
 for an explosive mixture? In what proportion should arsine and 
 oxygen be mixed? 
 
 4. Phosphorus forms an oxychloride, POC1 3 . How many 
 grams of phosphorus pentachloride and how many grams of water 
 would be required to prepare 100 grams of the compound? 
 
 5. How much blue vitriol, CuS04.5H 2 0, will be required to 
 give one pound (453 grams) of Paris green? 
 
CHAPTER XVII 
 GROUP IV: CARBON 
 
 Position of Carbon among the Elements. If we arrange 
 the elements -of low atomic weight in order, omitting hydro- 
 gen, we have the following list : 
 
 He, 4; Li, 7; Be, 9;B, 11; C, 12; N, 14; O, 16; F, 19; Ne, 20 
 
 In this list carbon stands half way between the two noble 
 gases helium and neon and also half way between the strongly 
 metallic element lithium and the most markedly non-metallic 
 element fluorine. In its properties carbon corresponds to 
 this unique position. In one of its elementary forms, the 
 diamond, it is transparent and a non-conductor of electricity . 
 In another form, graphite, it is opaque, has a metallic 
 luster and is a fairly good conductor of electricity. Car- 
 bon also shows an extraordinary power of combining with 
 itself and as a result of this power it forms an almost end- 
 less variety of complicated compounds. Many of these 
 are found in vegetable and animal substances and carbon is 
 the element of preeminent importance in all living bodies. 
 
 The ultimate source of all of the carbon for the growth of 
 trees and plants is the carbon dioxide of the air. Although 
 it forms only 0.03 per cent of the volume of the air, growing 
 plants secure their supply of carbon from this, absorbing 
 from sunlight the energy by means of which they seize upon 
 carbon and give the oxygen back to the atmosphere. 
 
 Diamonds. Diamonds are found in only a few places in 
 the world. The largest supply thus far discovered is at 
 Kimberly, in South Africa, and the company owning the 
 
 185 
 
186 CARBON 
 
 mines there now controls the diamond markets of the 
 world. The output of the mines is valued at $15,000,000 
 annually. 
 
 The diamond is the hardest substance known. It also 
 has a very high index of refraction and a very high dispersive 
 power; that is, a diamond prism separates red and violet 
 light widely in the spectrum. Because of these properties 
 diamonds which are cut in suitable forms reflect and re- 
 fract the light which falls on them and furnish the most 
 brilliant gems. 
 
 Until near the close of the nineteenth century no one was 
 able to discover how diamonds were formed in nature or 
 how they might be produced artificially. Finally, the French 
 chemist Moissan discovered some microscopic diamonds in a 
 meteor which fell in Canon Diablo, Texas. This led him to 
 the thought that he had, as it were, caught nature in the 
 act of making diamonds. After considering the conditions 
 under which the diamonds must have been formed in the 
 meteor he took some iron and melted it and heated it to a 
 very high temperature, above the melting point, in an elec- 
 tric furnace. He then dissolved some charcoal in the iron 
 and plunged the mass into water. This caused the rapid 
 solidification of the exterior while the interior portions were 
 still at a high temperature and liquid. Iron expands as it 
 solidifies and a considerable portion of the dissolved carbon 
 separates, ordinarily in the form of graphite. With the solid 
 shell on the outside the carbon separated from the iron 
 under conditions of high pressure. Moissan demonstrated 
 that some of the crystals which separated were diamonds, 
 by the following tests: 1. The crystals sank in a liquid in 
 which graphite will float. The specific gravity of the 
 diamond is 3.5; that of graphite is 2.25. 2. The crystals 
 scratched the face of a ruby. Only a diamond or the arti- 
 ficial mineral carborundum will do this and the specific 
 gravity of the crystals was greater than that of carborun- 
 
GRAPHITE 187 
 
 dum. 3. When the crystals were burned in a current of 
 oxygen 12 parts by weight gave 44 parts of carbon dioxide. 
 
 Diamonds are used in cutting glass but an edge produced 
 by cleavage must be employed. Black diamonds of an 
 inferior quality are used for drills. These are used espe- 
 cially for boring in such a manner that a solid core of rock 
 may be removed. By this means a solid piece of rock from 
 the bottom of a deep well may be brought to the surface 
 and examined. 
 
 Graphite. This allotropic variety of carbon is found 
 much more abundantly and it may also be readily prepared 
 by crystallizing carbon from melted cast iron and by heating 
 impure carbon to a very high temperature in an electric 
 furnace. Much of the graphite of commerce is now manu- 
 factured in this manner. 
 
 Graphite is sometimes called " black lead" and "lead" 
 pencils are made from it. It is, of course, wholly different 
 from metallic lead. The graphite for lead pencils is 
 specially prepared by treatment with chemicals, mixing 1 
 it with a little clay, and pressing it into blocks from which 
 the leads of different hardness and quality are to be cut. 
 
 Graphite is used as a lubricant, especially for wooden 
 bearings, for stove polish and in crucibles designed for use 
 at very high temperatures for melting steel and refractory 
 alloys. For crucibles the graphite is mixed with a small 
 amount of clay to bind the particles together. The graphite 
 on the surface of the crucible burns out and the clay then 
 protects the remainder. 
 
 Graphite does not melt even at the temperature of the 
 electric arc (about 3600) but it volatilizes rapidly. It 
 also sublimes slowly at the temperature of the carbon fila- 
 ment in an electric light bulb. 
 
 Amorphous Carbon. When almost any compound of 
 carbon or almost any animal or vegetable substance is 
 heated to a high temperature it turns black and a substance 
 
188 CARBON 
 
 which is called amorphous carbon separates. It is only 
 with considerable difficulty that by heating some substance, 
 such as sugar, which contains only carbon, hydrogen and 
 oxygen, a pure, amorphous carbon can be obtained. Char- 
 coal, lampblack, coke and all of the other forms of amor- 
 phous carbon contain hydrogen and are otherwise far from 
 pure carbon. 
 
 Amorphous carbon does not have sharply defined prop- 
 erties. The density varies from 1.45 up to that of graphite. 
 The kindling temperature of some forms is as low as 300 
 while that of others approaches that of graphite. 
 
 Lampblack is prepared by burning rosin, petroleum and 
 similar substances under conditions to produce a smoky 
 flame. It is used as a pigment, especially in printer's ink. 
 Its value depends in part on the fact that, in common with 
 other forms of carbon, there is nothing that will dissolve it 
 at ordinary temperatures and it is absolutely permanent 
 and unaffected by light. 
 
 Charcoal. When wood is heated, water, tar and com- 
 bustible gases are driven out and there finally remains 
 a black, porous residue retaining the form of the wood 
 and consisting chiefly of carbon with the mineral constit- 
 uents originally present in the wood. This residue is 
 called charcoal. It was formerly manufactured by piling 
 wood in heaps, covering it with turf and allowing it to burn 
 with a smoldering fire until the conversion to charcoal 
 was complete. This wasteful process has been largely 
 replaced by the use of apparatus for heating the wood in 
 closed retorts in such a manner as to save the valuable 
 products which distil away, as well as the charcoal. Large 
 quantities of charcoal were once used in the manufacture of 
 iron, and at one time the destruction of the forests in England 
 was threatened by this use. It has now been displaced by 
 coal and coke. It is still sometimes used by tinners and 
 it is used for the filtration of alcohol in its purification. 
 
COKE AND BY-PRODUCTS J89 
 
 Charcoal will condense many times its volume of am- 
 monia, hydrogen sulfide and other gases, especially those 
 gases which are easily liquefied. The gases probably con- 
 dense to the liquid form in the minute pores of the charcoal. 
 They are expelled by heat and freshly heated charcoal is 
 most effective for the absorption of the gases. 
 
 Different kinds of charcoal vary greatly in their absorp- 
 tive power and especially effective kinds made from coconut 
 shells have been developed for the gas masks used in war. 
 
 Bone-black and Animal Charcoal are prepared by char- 
 ring bones and slaughter-house refuse. They are used 
 especially to remove color from syrup in the purification 
 of sugar. 
 
 Coke is manufactured by heating bituminous coal. 
 The larger part of the coke used in the United States is 
 made in beehive coke ovens, round 
 brick chambers about 12 feet in 
 diameter and 7^ feet high. These 
 are partly filled with the coal and 
 the gases given off from the coal 
 are allowed to burn above the 
 surface of the coal within the 
 
 chamber (Fig. 34) . The burning gases furnish the heat to 
 coke the coal. 
 
 By-products Coke Ovens. In Germany, and more and 
 more in this country, this wasteful method is being replaced 
 by coking ovens. Flues are so constructed that the gas, tar 
 and ammonia water may be recovered. Nearly all of the 
 ammonia of commerce comes from this source. The 
 coal tar is the source from which benzene, toluene, naph- 
 thalene and other hydrocarbons used in the manufacture 
 of dyes and explosives are obtained. Large quantities of 
 the tar distillates are also used to " creosote" lumber for 
 railway ties, pavements and other purposes to preserve 
 it from decay. The gas produced is more than enough to 
 
190 
 
 CARBON 
 
 give the heat required for coking the coal and may be 
 used for illuminating gas or in gas engines for the generation 
 of power. Coke is chiefly used in the manufacture of iron. 
 The ovens (Fig. 35), are so arranged that the air for the 
 combustion of the gas used in heating the retorts comes up 
 through a chamber filled with a checker- work of bricks. 
 The air and gas come together and burn in the space 
 
 FIG. 35. 
 
 between the chambers filled with coal. The heated prod- 
 ucts of combustion then pass down through a second set 
 of chambers filled with bricks which absorb the heat that 
 would otherwise be wasted. After a time the current of 
 air is reversed, going up through the heated chamber and 
 down through the cooled one. Such an arrangement is 
 called a " regenerative furnace.' 7 
 
GAS CARBON. COAL 191 
 
 Gas Carbon. Carbon Electrodes. A very dense form 
 of carbon, approaching graphite in its properties, is pre- 
 pared by grinding anthracite coal, petroleum coke and other 
 kinds of amorphous carbon, mixing the powder with a little 
 coal tar or some petroleum product to make it cohere 
 and, after subjecting it to a high pressure, heating the mix- 
 ture to 1200-1400 for one or two days. The carbon be- 
 comes very dense and hard and is a good electrical conductor. 
 It is used for the carbons of arc lights and for electrodes in 
 the electrolytic manufacture of chlorine and sodium hydrox- 
 ide from brines. 
 
 Coal. During millions of years of geological time vast 
 stores of vegetable materials were accumulated in different 
 parts of the world from the growth of luxuriant, tropical 
 forests. The woody fiber of this material contained only 
 about 50 per cent of carbon with approximately 6 per 
 cent of hydrogen and more than 40 per cent of oxygen. 
 During the ages which have passed since the accumulation 
 of the material it has undergone a slow process of carboniza- 
 tion, by which the oxygen with some of the hydrogen and a 
 little carbon have escaped, with the result that the per cent 
 of carbon has increased while the per cent of oxygen has 
 decreased, but the per cent of hydrogen has not greatly 
 changed except in the final transformation to anthracite 
 coal. This has given the series of products known as peat, 
 lignite, bituminous coal and anthracite. The first three 
 give off volatile matters rich in hydrogen and carbon when 
 heated and for this reason burn with a smoky flame unless 
 special precautions are taken to secure a smokeless combus- 
 tion. Anthracite coal yields almost no volatile matters 
 when heated and burns with a smokeless flame. 
 
 The changes which have occurred in these transformations 
 from woody fiber to anthracite coal can be seen from the 
 following table : 
 
192 
 
 CARBON 
 
 CHANGES OF WOOD MATERIAL DURING GEOLOGICAL TIME' 
 
 
 
 | 
 
 
 
 Percentage composition 
 
 i 
 
 
 
 exclusive of moisture 
 
 
 Calorific 
 
 
 and ash 
 
 Percent- i Percent- 
 
 value; 
 
 Material 
 
 
 age of 1 age of 
 
 calories 
 
 
 
 
 
 
 ash ' i moisture 
 
 per kilo- 
 
 
 Car- 
 
 Hydro- 
 
 Oxy- 
 
 Nitro- 
 
 
 gram 
 
 
 bon 
 
 gen 
 
 gen 
 
 gen 
 
 1 
 
 
 
 
 
 
 
 | | 
 
 Wood Oak.. . . 
 
 50.35 
 
 6.04 
 
 43.52 
 
 0.09 
 
 0.37 
 
 20. OO 2 
 
 3696 
 
 Peat 
 
 59.70 
 
 5.70 
 
 33.04 
 
 1.56 
 
 11.84 
 
 14.242 
 
 3979 
 
 Brown Lignite, 
 
 
 
 
 
 
 
 North Dakota 
 
 74.88 
 
 4.99 
 
 19.12 
 
 1.01 
 
 9.35 i 35.38 
 
 3846 
 
 Black Lignite, 
 
 
 
 
 
 | 
 
 Colorado 
 
 76.83 
 
 5.34 
 
 16.29 
 
 1.54 
 
 5.99 ! 18.68 
 
 5635 
 
 Bituminous, 
 
 
 
 
 
 
 
 Illinois. ....... 
 
 83.42 
 
 5.29 
 
 9.52 
 
 1.77 
 
 11.28 1 8.50 
 
 6542 
 
 Semibitumin- 
 
 
 
 
 
 
 
 
 ous, West Vir- 
 
 
 
 
 
 
 
 
 ginia Poca- 
 
 
 
 
 
 
 
 
 hontas 
 
 91.50 
 
 4.38 
 
 3.07 
 
 1.05 
 
 6.55 
 
 3.67 
 
 7939 
 
 Anthracite 
 
 93.76 
 
 2.72 
 
 3.11 
 
 0.41 
 
 10.80 
 
 2.18 
 
 7216 
 
 Charcoal 
 
 84.11 
 
 1.53 
 
 14.36 
 
 
 2.50 
 
 
 6626 
 
 Coke 
 
 95.47 
 
 0.67 
 
 2.82 
 
 1.04 
 
 14.80 
 
 
 6768 
 
 1 Table prepared by Professor S. W. Parr. 2 Air dry. 
 
 There are three kinds of bituminous coals: coking coals, 
 which sinter together when heated, giving a hard, coherent 
 coke suitable for use in blast furnaces for the manufacture 
 of cast iron; non-coking coals, which do not sinter, or sinter 
 imperfectly, giving a friable coke ; and cannel coals, having a 
 peculiar homogeneous structure with a conchoidal fracture. 
 The last burn with a brilliant, luminous flame and are used 
 in the manufacture of illuminating gas. 
 
 SUMMARY 
 
 Carbon is at the center of the first period of elements both 
 in position and properties. It is the most important element 
 in living bodies. 
 
 Plants obtain carbon from the carbon dioxide of the 
 atmosphere. 
 
 Diamonds are a very hard, transparent, dense form of car- 
 bon, valuable for gems, cutting glass and drills. 
 
EXERCISES; CARBON 193 
 
 Diamonds have been prepared artificially by crystallizing 
 carbon from iron under a high pressure. 
 , Graphite is found in nature and is prepared by crystalliz- 
 ing carbon from a solution in iron or by heating it in an 
 electric furnace. 
 
 Graphite is used in lead pencils, stove polish, crucibles 
 and as a lubricant. 
 
 Amorphous, or uncrystallized carbon is found in an im- 
 pure form in lampblack, charcoal, coke and coal. 
 
 Lampblack is used as a pigment and in printer's ink. 
 
 Charcoal is made by heating wood. It is used for small 
 fires, in filtering alcohol and to absorb noxious gases. 
 
 Bone-black and animal charcoal are used in purifying 
 sugar. 
 
 Coke is made by heating coal in beehive or in by-product 
 coke ovens. It is used in the manufacture of iron. 
 
 Gas carbon is used for the carbon electrodes of arc lights. 
 
 Vegetable matter has been transformed slowly into peat, 
 lignite, bituminous coal and anthracite coal, and in some 
 cases to graphite. 
 
 Bituminous coals are distinguished as coking, non-coking 
 and cannel coals. 
 
 EXERCISES 
 
 1. The heat of combustion of a gram atom of diamonds is 
 94,310 small calories. That of a gram atom of amorphous carbon 
 is 97,350 calories. Is the formation of diamonds from amorphous 
 carbon exothermic or endothermic? Will the formation be fav- 
 ored by an increase or decrease of the temperature (see p. 121)? 
 Will it be favored by pressure? 
 
 2. What volume of carbon dioxide will be formed by burning a 
 gram of carbon? What volume of air will be required to furnish 
 the oxygen for the combustion? 
 
 3. What volume of air will be required to burn a pound of 
 carbon? What weight of air? 
 
 13 
 
194 CARBON 
 
 4. What volume and weight of air are required to burn a pound 
 of hydrogen? 
 
 5. What volume and weight of air will be required to burn a 
 pound of bituminous coal containing 73 per cent of carbon and 
 4.7 per cent of combustible hydrogen? 
 
 6. How many kilograms of water ought, theoretically, to be 
 evaporated by the heat from burning a kilogram of coal which 
 gives by its combustion 6500 calories per kilogram? How many 
 pounds of water per pound of coal? How many pounds of water 
 evaporated by a pound of coal is considered good boiler efficiency? 
 
CHAPTER XVIII 
 HYDROCARBONS, GAS, FLAME 
 
 Contrast of Carbon with Other Elements. The elements 
 of the halogen family each form only a single compound 
 with hydrogen, of which hydrochloric acid, HC1, may be 
 taken as typical. Oxygen forms two compounds, water, 
 H 2 O, and hydrogen peroxide, H 2 02. Nitrogen forms four 
 or five compounds but only ammonia, NH 3 , is met with 
 in common experience. In contrast with these, carbon 
 combines with hydrogen to form several thousands of com- 
 pounds. No other element gives more than a very small 
 number of compounds with hydrogen. The compounds of 
 hydrogen with carbon are called hydrocarbons. The large 
 number of these is due to the power which carbon has of 
 combining with itself in a great variety of ways. 
 
 Series of Hydrocarbons. The compounds of carbon and 
 hydrogen are classified by arranging them in a series in 
 accordance with their formulas and structure. The series 
 which contains most hydrogen in proportion to the carbon is 
 called the marsh-gas series from the first member, methane 
 or marsh gas, CH 4 . Some members of the series are: 
 
 Methane, CH 4 
 
 Ethane, C 2 H 6 
 
 Propane, C 3 Hs 
 
 Butane, C 4 Hi 
 
 Pentane, C 5 Hi 2 
 
 Hexane, CeHu 
 
 Heptane, CrHie 
 
 Octane, CgHis 
 
 Nonane, CgH 2 o 
 
 Decane, CioH 2 2 
 
 195 
 
196 HYDROCARBONS, GAS, FLAME 
 
 The formula C n H 2n +2 may be used for any member of the 
 series, n standing for the number of carbon atoms. 
 
 Structure of the Hydrocarbons of the Methane Series. 
 It will be noticed that in the successive members of this 
 series each new carbon atom carries with it two new hydro- 
 gen atoms. In other words, there is a difference of one 
 carbon atom and two hydrogen atoms between each hydro- 
 carbon of the series and the one preceding or following it. 
 The same relation is found in all other series of hydro- 
 carbons. This is most easily explained by the use of two 
 well-established principles: 1. The valence of carbon is 
 almost always four. 2. Carbon atoms readily unite with 
 each other as well as with other elements. 
 
 The application of these principles leads to the following 
 structural formulas for the first four members of the series : 
 
 H H H H H H 
 
 II III 
 
 H-C-H H-C-C-H H-C-C-C-H 
 
 I 11 111 
 
 H H H H H H 
 
 Methane Ethane Propane 
 
 H H H H H 
 
 1 I I I I 
 
 H-C-C-C-C-H H-C-H 
 
 H 
 
 H 
 
 H H H H 
 
 Butane TI p O O TT 
 
 Boiling point +1 ' ^ " " y " " V ' 
 
 H H H 
 
 Isobutane 
 Boiling point 11.5 
 
 The formulas for butane and isobutane represent two 
 different compounds, C 4 Hi , which are actually known. 
 It would carry us too far to explain here how the two com- 
 pounds are distinguished. 
 
 Methane, CH 4 . Fire-Damp. Natural Gas. The first 
 member of the series is often called marsh gas. If the decay- 
 
DA\ Y SAFETY LAMP 
 
 197 
 
 FIG. 36. 
 
 ing leaves in the bottom of a pond are stirred up, an inflam- 
 mable gas, consisting largely of methane, usually comes up 
 to the surface and may be collected without difficulty. 
 The gas called fire-damp, which often causes 
 disastrous explosions in coal mines, consists 
 largely of methane. Natural gas also in 
 mostly methane. 
 
 Methane is most easily prepared in the 
 laboratory by heating a mixture of sodium 
 acetate and soda lime. The latter is a 
 mixture of sodium hydroxide, NaOH, and 
 slaked lime, Ca(OH) 2 . 
 
 Davy Safety Lamp. Mixtures of air and 
 methane in the right proportion explode 
 violently. A particularly distressing acci- 
 dent in a coal mine from this cause, early 
 in the nineteenth century, led to a request 
 that Sir Humphrey Davy should investigate the cause of 
 the explosion and, if possible, suggest a remedy. He 
 established, first of all, that when marsh gas is mixed with 
 from six to fourteen times its volume 
 of air the mixture may explode 
 violently. He also found that a com- 
 paratively high temperature, ap- 
 proaching a red heat, is required to 
 ignite the mixture. He then invented 
 a lamp in which the flame is com- 
 pletely surrounded with heavy wire 
 gauze. With such a lamp, Fig. 36, 
 the mixture of gases on the outside of 
 the gauze does not become heated to 
 its kindling temperature and so does 
 not explode. Precautions are always taken, however, to 
 avoid the presence of explosive mixtures in mines. Mix- 
 tures of coal dust, flour dust, cr other fine particles of organic 
 
 FIG. 37. 
 
198 HYDROCARBONS, GAS, FLAME 
 
 matter, with air, may also explode violently. An explosion 
 of flour dust in a mill in Minneapolis once caused the de- 
 struction of the mill and the loss of several lives. Since 
 then great care is taken to prevent the accumulation of 
 dust in flour mills and factories. The effect of wire gauze 
 in cooling a gas below its kindling temperature is shown 
 in Fig. 37. 
 
 Petroleum. Methane, ethane, propane and butane, the 
 first four members of the methane series, are gases at ordi- 
 nary temperatures but the higher members of the series are 
 liquids or solids. A number of these are found in petroleum, 
 the oil which is obtained in many places by boring wells in 
 the earth. While petroleum has been known for a very 
 long time its use was first developed on a large commercial 
 scale in Pennsylvania in 1859. Since then oil has been 
 found in very many different places both in the United 
 States and in foreign countries. Ohio, Indiana, Illinois, 
 Texas, Oklahoma, Kansas and California may be men- 
 tioned as states producing large quantities of petroleum. 
 Canada, Mexico and the region of the Caucasus may be 
 mentioned among foreign countries. There is reason to 
 think that immense fields of petroleum remain undiscovered. 
 
 Crude petroleum is often used as a fuel. It is purified 
 for the production of kerosene, gasoline and other products, 
 chiefly by distillation, partly by treatment with concen- 
 trated sulfuric acid to remove substances which give the 
 oi! a disagreeable odor or interfere with its proper burning. 
 
 Gasoline consists of the lower boiling constituents which 
 may be converted into a vapor at a sufficiently low tempera- 
 ture to form an explosive mixture with air. The use of 
 gasoline in engines, especially in automobiles, is familiar to 
 everyone. The successful use in an engine depends very 
 largely on securing a proper mixture of air with gasoline 
 vapor. If too much gasoline is used, a part burns only to 
 carbon monoxide and there is a great loss of energy. Under 
 
UNSATURATED COMPOUNDS 
 
 some conditions a part of the carbon fails to burn at all 
 and is deposited in the cylinders of the engine. The carbon 
 deposited in this way often interferes seriously with the 
 operation of the engine. A mixture of air with too much 
 gasoline may fail to explode when the engine is hot, just 
 as one with too little vapor fails to explode when the engine 
 is cold. 
 
 Kerosene, which is used in lamps and stoves, should 
 be free from substances which are volatile enough to give 
 an explosive mixture with air at moderate temperatures. 
 In most states the law requires that the flashing point of 
 kerosene shall be above 150 F., i.e., so much vapor that it 
 can be ignited with a flame must not be given off below 
 that temperature. With the increasing demand for gaso- 
 line, manufacturers of kerosene are no longer tempted to 
 sell kerosene with a low flashing point. 
 
 Vaseline, 1 paraffin and lubricating oils are the other best 
 known products obtained from petroleum. 
 
 Ethylene, C 2 H 4 . Unsaturated Compounds. When alco- 
 hol, C 2 H 6 O, is heated with concentrated sulfuric acid the 
 elements of water are removed and a gas, ethylene, C 2 H 4 , is 
 produced. It may also be prepared almost quantitatively, 
 by passing alcohol over aluminium oxide heated to a mod- 
 erate temperature. 
 
 Ethylene is also a constituent of coal gas and of other 
 illuminating gases and is one of the most important of the 
 compounds which give luminous quality to such gases. 
 
 Ethylene combines directly with chlorine or bromine to 
 form such compounds as ethylene chloride, C 2 H 4 C1 2 , and 
 ethylene bromide, C 2 H 4 Br 2 . Because of this property it is 
 said to be unsaturated. Hydrocarbons of the methane 
 series, which do not act in the same way but which give, 
 instead, substitution products, such as CH 3 C1 and C 2 H 5 C1, 
 
 1 Vaseline is a proprietary name used by the Chesebrough Manufac- 
 turing Company. The name used in the Pharmacopoeia and by other 
 manufacturers is petrolatum. 
 
200 HYDROCARBONS, GAS, FLAME 
 
 are said to be saturated. The difference will be clearer 
 from the following structural formulas: 
 
 H 
 
 I Cl 
 
 H-C-H | 
 
 H-C -H 
 H-C-H + C1-C1= | 
 
 | H-C-H + H-C1 
 
 H | 
 
 Ethane JJ 
 
 Ethyl chloride 
 
 Cl 
 
 H-C-H H-C-H 
 
 + C1-C1= | 
 
 H-C-H H-C-H 
 
 Ethylene 
 
 Cl 
 
 Ethylene dichloride 
 
 Acetylene, C 2 H 2 . Calcium carbide, CaC 2 , is now made 
 on a large scale by heating a mixture of lime and coke in an 
 electric furnace: 
 
 CaO + 3C = CaC 2 + CO 
 
 Calcium Carbon 
 carbide monoxide 
 
 When calcium carbide is put in water it reacts with it, 
 forming acetylene, C 2 H 2 , and calcium hydroxide: 
 
 CaC 2 + 2HOH = Ca(OH) 2 + C 2 H 2 
 
 Acetylene is a colorless gas which burns, under proper 
 conditions, with a brilliant, luminous flame. A cubic foot 
 of the gas may be made to give more than ten times as much 
 light as a cubic foot of good illuminating gas. 
 
 The luminous quality of the acetylene gas flame is due 
 to the separation of particles of carbon which are heated to 
 a very high temperature in the flame. A special burner, 
 which will secure complete combustion of the carbon with- 
 out the escape of smoke, is required. 
 
ENDOTHERMIC COMPOUNDS 201 
 
 Various forms of generators for acetylene are in use. The 
 forms in which the carbide is dropped into water are more 
 satisfactory than those in which water is dropped on the 
 carbide. In the latter forms so much heat is generated 
 that a part of the acetylene polymerizes and is lost. 
 
 Endothermic Compounds. When acetylene is heated it 
 decomposes into carbon and hydrogen: 
 
 O2ll2 = 2O "f" Il2 
 
 Considerable heat is evolved during the decomposition. 
 This has been proved by comparing the heat of combustion 
 of acetylene with that of amorphous carbon and hydrogen. 
 
 26 g. of acetylene give when burned 313.8 calories 
 
 24 g. of carbon give when burned . .195.3 
 2 g . of hydrogen give when burned . 68 . 4 263 . 7 
 
 Difference 50 . 1 
 
 These results show that when 24 grams of carbon combine 
 with 2 grams of hydrogen to form acetylene 50 calories of 
 heat are absorbed, or that when 26 grams of acetylene 
 decompose into carbon and hydrogen 50 calories will be 
 given off. 
 
 These facts are of practical importance from two points 
 of view: 
 
 1. The solid carbon which separates when acetylene is 
 heated is the source of the intense light of the acetylene 
 flame (see above). 
 
 2. If the decomposition of liquid acetylene is once 
 started, the heat generated by the decomposition hastens 
 it and the large volume of hydrogen formed may cause an 
 explosion. It is, in fact, possible to explode liquid acetylene 
 with a cap, very much as nitroglycerine is exploded. When 
 these properties of acetylene became generally known laws 
 were passed forbidding the use of liquid acetylene. It has 
 been discovered, however, that acetone or acetaldehyde 
 
202 HYDROCARBONS, GAS, FLAME 
 
 will absorb large quantities of acetylene under pressure 
 and give it out again when the pressure is released and that 
 the solution, if not too rich in acetylene, is not explosive. 
 Such a solution is now used, especially for automobile lights. 
 
 Such compounds as acetylene, which absorb heat in 
 their formation, are called endothermic compounds and 
 the reactions by which they are prepared are called endo- 
 thermic reactions. 
 
 Reactions of this type are favored by a high tempera- 
 ture (p. 121), and in accordance with this acetylene is 
 formed by the direct union of carbon and hydrogen at 
 the temperature of the electric arc. 
 
 Benzene, C 6 H 6 . If acetylene is passed through a tube 
 heated to a moderate temperature a part of it will poly- 
 merize to benzene: 
 
 Benzene and several other closely related hydrocarbons 
 are found in coal tar. They are obtained from the tar by 
 distillation. These compounds are used in large quantities 
 for the manufacture of dyes, phenol ("carbolic acid") 
 and many medicinal products. 
 
 Toluene, C 7 H 8 , the second compound of the benzene 
 series, is used in making trinitrotoluene ("T. N. T."), 
 one of the explosives used in the Great War. 
 
 Other Hydrocarbons. Each of the hydrocarbons men- 
 tioned is the first member of a series. Other hydrocarbons 
 of the ethylene series have the formulas C 3 H 6 , C 4 H 8 , 
 C 5 Hi , C 6 Hi 2 , etc.; those of the acetylene series are C 3 H 4 , 
 C 4 H 6 , C 5 H 8 , C 6 Hio, etc., and those of the benzene series 
 are C 7 H 8 , C 8 Hi , C 9 Hi 2 , Ci Hi 4 , etc. Besides these 
 hydrocarbons and very many others belonging to these 
 four series there are many other series and several thousand 
 compounds containing only carbon and hydrogen are known. 
 
 Illuminating Gas. A little more than a century ago 
 
WATER GAS 203 
 
 the only forms of illumination in use were candles and 
 lamps of a comparatively inferior type, using lard or some 
 kind of vegetable or animal oil. Early in the nineteenth 
 century the manufacture of an illuminating gas by heating 
 bituminous or cannel coal was gradually introduced in 
 large cities. The gas formed in this way is a complex 
 mixture of hydrogen, methane, ethylene, acetylene, ben- 
 zene vapor and other hydrocarbons, with small quantities 
 of carbon monoxide, carbon dioxide and hydrogen sulfide. 
 The hydrogen sulfide is mostly removed before the gas 
 is used. 
 
 When gas is burned in a flat flame or in an Argand burner 
 the illuminating quality depends largely on the so-called 
 " heavy hydrocarbons," ethylene, acetylene, benzene, etc., 
 which are present. If burned with a Welsbach mantle 
 the light given is nearly proportional to the heat of com- 
 bustion of the gas, but the heavy hydrocarbons give much 
 more heat in proportion to their volume than the other 
 constituents do and are still important in determining the 
 quality of the gas. 
 
 Water Gas. For the manufacture of water gas a suitable 
 furnace is filled with a large mass of coke and this is brought 
 to a high temperature by burning some of it with a blast 
 of air. The air is then cut off and steam is turned into 
 the chamber for a few minutes. This reacts with the coke 
 in accordance with the equation: 
 
 C + H 2 = CO + H 2 
 
 This reaction is an endothermic one as is seen from the 
 following relations : 
 
 12 grams of amorphous carbon give 97 . 65 calories 
 
 28 grams of carbon monoxide give 68. 2 
 
 2 grams of hydrogen (burned to 
 
 steam) give 58.0 
 
 Excess 28 . 55 calorie; 
 
 126.2 
 
204 HYDROCARBONS, GAS, FLAME 
 
 In order, therefore, to convert 12 grams of carbon and 
 18 grams of steam into 28 grams of carbon monoxide and 
 2 grams of hydrogen, 28.5 calories of heat must be added 
 from some source. Evidently the only way in which this 
 heat can be furnished is by the hot coke and that must cool 
 off rapidly as the reaction proceeds. After a few minutes it 
 is necessary to turn off the steam and turn on the blast of 
 air to bring the coke again to a high temperature. 
 
 While the steam is passing in, the mixture of carbon 
 monoxide and hydrogen formed is saved and is known as 
 water gas. Large volumes of this gas can be manufactured 
 very cheaply by the process. 
 
 Pure water gas burns with the blue flame characteristic 
 of carbon monoxide and gives almost no light. The heat 
 of combustion, too, is only about one-half that of a good 
 illuminating gas. For use in illumination it is enriched by 
 adding to it the mixture of rich gases obtained by heating 
 petroleum. 
 
 The most serious objection to the use of water gas in 
 illuminating gas is the poisonous character of the carbon 
 monoxide which it contains. This has led some States to 
 prohibit its use. There is additional danger from the fact 
 that water gas has much less odor than coal gas, and leaks 
 are not so quickly noticed. 
 
 Producer Gas. It would be possible, theoretically, to 
 pass a mixture of air and steam over a mass of coke or 
 coal in such proportions that the burning of part of the 
 coal by the air would furnish enough heat to just keep up 
 the temperature of the mass and still convert a part of 
 the steam and coal into water gas. The gas made in this 
 manner would give just as much heat in its combustion 
 as could be obtained by burning the original coal. Of 
 course some heat, must be lost in carrying out such a 
 scheme practically, but apparatus has been devised by 
 which coal may be converted into combustible gases that 
 
CANDLE FLAME 205 
 
 retain 80 to 85 per cent of the original heating power of 
 the coal. Such a gas is called " producer gas." It consists 
 chiefly of a mixture of hydrogen, carbon monoxide, CO, 
 and nitrogen. 
 
 Producer gas can be used much more economically than 
 the original solid fuel in some processes for the manu- 
 facture of steel and glass, and in gas engines. 
 
 Flames. With the exception of the mercury vapor elec- 
 tric lamp, all forms of illumination in common use depend 
 on heating some solid to a high temperature. 
 
 All of the older forms of illumination depend on the 
 combustion of gaseous compounds of carbon under such 
 conditions that solid particles of carbon exist momentarily 
 in the flame. In the candle or lamp the material burned 
 is converted into a gas immediately before combustion 
 but such flames are gaseous flames just as much as a flame 
 of illuminating gas. 
 
 Candle Flame. In the candle flame we may distinguish 
 three parts: (1) An inner portion of unburned gas, sur- 
 rounding the wick. This is relatively cool, as will be seen 
 on holding a wooden toothpick or match stick across the 
 flame. It will char at the edges of the flame before it does 
 in the center. (2) A luminous zone of partial combustion. 
 Here the carbon, which separates in the decomposition 
 of compounds present, is heated red hot and gives the 
 light of the flame. The carbon is deposited on any cold 
 object held in the flame. (3) A zone of complete combus- 
 tion surrounding the flame and most apparent at the base 
 of the flame. Here the carbon and hydrogen of the flame 
 are completely burned to carbon dioxide and water. 
 
 Bunsen Burner. In the Bunsen burner, Fig. 37 (p. 197), 
 the gas is mixed with an amount of air insufficient for 
 complete combustion before it is ignited at the top of the- 
 burner. This causes the carbon to burn to carbon monox- 
 ide with little or no separation of carbon. The gases burn 
 
206 
 
 HYDROCARBONS, GAS, FLAME 
 
 in the inner cone of the flame, giving a mixture of carbon 
 monoxide, hydrogen, carbon dioxide, water vapor and 
 nitrogen. In the outer zone of the flame the carbon monox- 
 ide and hydrogen burn to carbon dioxide and water vapor. 
 
 27201 
 
 Fio. 38. 
 
 FIG. 39. 
 
 Temperature of Flames, The temperature of the flame 
 'of a Bunsen burner varies from about 300 in the center, 
 near the mouth of the burner, where combustion has not 
 begun, to about 1550 in the portion between the inner cone 
 and the outside of the flame. These temperatures are 
 shown in detail in Fig. 39. 
 
BLOWPIPE 207 
 
 In the Meker burner, Fig. 38, by widening the top of the 
 burner and giving it a considerable number of fairly heavy 
 metallic partitions, the inner cone is divided into a number 
 of small and very short parts. This brings the high tem- 
 perature of the upper part of the Bunsen flame down close 
 to the mouth of the burner, concentrates the flame and 
 gives it a more uniform and somewhat higher temperature. 
 
 The temperatures given in the figures are, of course, the 
 temperatures of the flame when no substance radiating heat 
 is present. A platinum or porcelain crucible placed in the 
 flame will be at a much lower temperature. A 20-gram 
 platinum crucible placed 1 cm. above the Meker burner, 
 with ordinary gas, will usually have a temperature of- 
 900-950. 
 
 Blowpipe. By means of a blowpipe, Fig. 40, the flame of 
 a candle, or the luminous flame burning at the slanting end 
 of a tube introduced in a 
 Bunsen burner, may be 
 changed to a narrow, 
 pointed flame which can 
 be used to advantage for 
 
 either the oxidation or reduction of substances laid on a 
 stick of charcoal or dissolved in a bead of borax glass. The 
 extreme tip of the flame gives an oxidizing effect, because it is 
 intensely hot and the oxygen of the air can act on the sub- 
 stance. The inner, faintly luminous cone of the flame is 
 reducing in its action, owing to the unburned, combustible 
 gases present. 
 
 SUMMARY 
 
 Carbon, in contrast with other elements, forms several 
 thousands of compounds with hydrogen. 
 
 Because of the valences of carbon and hydrogen, the 
 successive members of each series of hydrocarbons differ 
 by one carbon and two hydrogen atoms. 
 
208 HYDROCARBONS, GAS, FLAME 
 
 Methane is found in natural gas and in fire-damp. It is 
 prepared by heating sodium acetate with soda lime. 
 
 Methane and other hydrocarbons give explosive mix- 
 tures with air. Flour dust or coal dust may also explode 
 when mixed with air and ignited. 
 
 The Davy safety lamp was invented to prevent explosions 
 in coal mines. 
 
 Gasoline, kerosene, lubricating oils, vaseline and paraffin 
 are the most important products obtained from petroleum. 
 
 The flashing point of gasoline should be low, for satis- 
 factory use, that of kerosene should be moderately high, 
 for safety. 
 
 Ethylene is prepared by the decomposition of alcohol. 
 It gives a luminous flame. It is unsaturated and combines 
 directly with chlorine or bromine. 
 
 Calcium carbide is prepared by heating lime and carbon 
 in an electric furnace. 
 
 Acetylene is prepared by the action of water on calcium 
 carbide. It is an endothermic compound and decomposes 
 into its elements with an evolution of heat. Liquid acety- 
 lene is explosive. 
 
 Benzene and many other hydrocarbons are obtained from 
 coal tar. These are used in making dyes, carbolic acid or 
 phenol, explosives and many medicinal products. 
 
 Illuminating gas is manufactured from coal or by enrich- 
 ing water gas. 
 
 Water gas is made by passing steam over hot coke. 
 The reaction is endothermic and intermittent for that 
 reason. Water gas is poisonous. 
 
 Producer gas is made by burning coal with a limited 
 supply of air. 
 
 The luminosity of candle flames and of gas flames is due 
 to heated particles of carbon. 
 
 In the Bunsen burner the introduction of air at the base 
 of the burner causes the partial combustion of the gas 
 
EXERCISES; HYDROCARBONS 
 
 209 
 
 within the flame and this prevents the separation of 
 carbon. 
 
 A blowpipe may be used to give either an oxidizing or 
 a reducing flame. 
 
 EXERCISES 
 
 1. What are the general formulas of the hydrocarbons of the 
 ethylene, acetylene and benzene series corresponding to the 
 formula C w H 2n+2 for the methane series? 
 
 2. What is the general formula of the hydrocarbons of a series 
 intermediate between the acetylene and benzene series? 
 
 3. Write the equations for the combustions of the following 
 hydrocarbons: methane, ethane, acetylene, benzene. In what 
 proportion by volume must each gas be mixed with oxygen for 
 complete combustion? In what proportion with air? 
 
 4. If gasoline has an average composition corresponding to the 
 formula C 6 Hi 4 , in what proportion by volume must its vapor be 
 mixed with air for complete combustion? 
 
 5. What weight of gasoline will be required to form the most 
 effective explosive mixture with one cubic foot (28.3 liters) of 
 air? 
 
 6. Find the dimensions of the cylinder of an automobile engine 
 and calculate how much gasoline should be introduced for one 
 stroke of the piston. 
 
 7. The heats of combustion of one cubic foot of the more im- 
 portant constituents of illuminating gas are as follows in calories 
 and in British Thermal Units: 
 
 
 Heat of t 
 
 ombustion 1 
 
 
 Calories 
 
 B. T. U. 
 
 Carbon monoxide, CO. 
 
 77 4 
 
 307 
 
 Hydrogen, Ha 
 
 66 3 
 
 263 
 
 Methane, CH 4 
 
 215 
 
 853 
 
 Ethylene, C 2 H 4 
 
 357 6 
 
 1420 
 
 
 
 
 Kiases at 15.6 (60 F.) burned to vapors at 164 (328 F.). 
 =3.968 British Thermal Units. 
 14 
 
 1 calorie 
 
210 
 
 HYDROCARBONS, GAS, FLAME 
 
 The products of combustion are assumed to be carbon dioxide 
 and liquid water. 
 
 Assuming that the "heavy hydrocarbons" consist essentially 
 of ethylene, what will be the heat of combustion of one cubic foot 
 of the following samples of gas: 
 
 
 Coal gas 
 
 Enriched 
 water gas 
 
 Producer gas 
 
 Carbon dioxide, C02. 
 
 1 1 
 
 3 
 
 1 5 
 
 Carbon monoxide, CO . 
 
 7 2 
 
 26 1 
 
 23 5 
 
 Hydrogen, HS 
 
 49 
 
 32 1 
 
 6 
 
 Methane, CH 4 
 
 34.5 
 
 19.8 
 
 3.0 
 
 Heavy hydrocarbons. 
 
 5 
 
 16 6 
 
 
 Nitrogen 
 
 3.2 
 
 2.4 
 
 66.0 
 
 
 
 
 
 Candle-power 
 
 100.0 
 17.5 
 
 100.0 
 25.0 
 
 100.0 
 
 
 
 
 
 8. The heat of combustion of one gram atom of carbon is 
 97,650 calories (small). The heat of combustion of one gram 
 molecule of hydrogen burned to liquid water is 68,400 calories. 
 The heat of combustion of one gram molecule of methane, CH 4 , 
 burned to carbon dioxide and liquid water is 214,000 calories. Is 
 the reaction expressed by the equation: 
 
 C + 2H 2 =* CH 4 
 
 exothermic or endothermic? Will the equilibrium be shifted 
 toward the formation of methane by a low or a high temperature? 
 
CHAPTER XIX 
 
 CARBON MONOXIDE, CARBON DIOXIDE, CARBON 
 DISULFIDE, CYANIDES 
 
 Formation of Carbon Monoxide. Whenever fuels burn 
 in a thick layer and especially in the burning of anthracite 
 coal, blue flames will be seen playing over the surface. 
 These are flames of carbon monoxide burning to carbon 
 dioxide. In the lower part of the mass of fuel some of it 
 may be burned to carbon dioxide and the latter is reduced 
 to the monoxide as it passes through the hot coals above: 
 
 CO 2 + C = 2CO 
 
 The reaction is an endothermic one and only occurs at a 
 high temperature. Twelve grams of carbon burned to 
 carbon monoxide give 29.45 calories (large), while the same 
 weight of carbon burned to carbon dioxide gives 97.65 
 calories. From this relation it will be seen that conditions 
 of combustion in a furnace or in a gas engine which lead to 
 the escape of carbon monoxide are very wasteful of the 
 energy of the fuel. It is partly for the same reason that it 
 is possible to convert coal into a producer gas which retains 
 a large per cent of its original energy (p. 204). 
 
 Preparation and Properties of Carbon Monoxide. 
 Carbon monoxide is most easily prepared in the laboratory 
 by heating a mixture of oxalic acid and concentrated sul- 
 furic acid: 
 
 H 2 C 2 O 4 + H 2 S0 4 = CO + C0 2 + H 2 SO 4 .H 2 O 
 
 Oxalic acid 
 
 The sulfuric acid catalyzes the decomposition of the 
 oxalic acid and combines with the water formed. The 
 
 211 
 
212 CARBON OXIDES, CYANIDES 
 
 carbon dioxide formed at the same time may be removed by 
 passing the gas through a tube filled with soda lime. 
 
 Carbon monoxide is a colorless, odorless gas. It burns 
 with a characteristic blue flame that gives very little 
 light. 
 
 Carbon monoxide is a very dangerous poison. Air con- 
 taining one part in a thousand of the gas is unsafe to breathe, 
 even for a short time. Air containing smaller amounts 
 may cause serious injury in the air of living rooms. The gas 
 seems to combine with the hemoglobin of the blood and to 
 so change it that it is no longer capable of performing its 
 normal function of transferring the oxygen of the air to the 
 tissues of the body. For this reason it acts as a cumulative 
 poison and recovery from poisoning with the gas is very 
 slow. This property of carbon monoxide constitutes a very 
 serious objection to the use of water gas. 
 
 The escape of carbon monoxide from a charcoal fire and 
 from a flame burning against a cold surface is the reason 
 why the products of combustion of neither should be per- 
 mitted to escape into a living room. 
 
 Carbon Dioxide. The formation of carbon dioxide by the 
 burning of carbon and its compounds and by respiration 
 have been repeatedly referred to. The gas is most easily 
 prepared by the treatment of a carbonate with hydrochloric 
 or sulfuric acid. Carbonic acid, H 2 CO3, very readily de- 
 composes into its anhydride, CO 2 , and water, and as carbon 
 dioxide is not very soluble in water the escape of the carbon 
 dioxide causes the reactions to go nearly to completion in 
 the direction of its formation : 
 
 CaCO 3 + 2HC1 <= CaCl 2 4- H 2 CO 3 
 
 Calcium 
 carbonate 
 
 H 2 CO 3 <= CO 2 + H 2 O 
 
 Carbon dioxide is a colorless gas about one-half heavier 
 than air. It is not distinctly poisonous when present in 
 
HENRY'S LAW 213 
 
 moderate amounts. It may, however, cause death from 
 suffocation. The gas sometimes accumulates in wells and 
 mines and is known as choke-damp. "Where its presence 
 is suspected the place should be tested with a burning 
 candle before it is entered, though it is possible to live for a 
 short time in a room containing so much carbon dioxide that 
 a candle is extinguished. 
 
 Carbon Dioxide and Water. Henry's Law. Soda Water. 
 -At ordinary temperature and atmospheric pressure, 
 water dissolves about its own volume of carbon dioxide. 
 From air containing carbon dioxide the gas will be dissolved 
 in proportion to the " partial pressure" of the gas present. 
 Thus the "partial pressure" of carbon dioxide in air con- 
 taining 10 per cent of the gas by volume is one-tenth of an 
 atmosphere and only one-tenth as much carbon dioxide will 
 be absorbed from such a mixture as will be absorbed if 
 water were in contact with the pure gas. This is known as 
 Henry's Law and applies to all gases that are slightly 
 soluble in water. 
 
 Under pressure water still takes up its own volume of 
 the gas but this means, of course, that the weight of the 
 gas absorbed increases proportionally with the pressure. 
 When the pressure is removed the excess of gas escapes with 
 effervescence. This property is used in soda water and in 
 carbonated beverages. 
 
 The solution of carbon dioxide in water has a faintly 
 acid taste and reddens blue litmus. The taste is only 
 slightly acid because the ionization : 
 
 - H 2 CO 3 <=* H+ + HCO 3 - 
 
 is only very slight, most of the carbon dioxide remaining 
 either as carbonic acid,H 2 C0 3 ,or, perhaps, as carbon dioxide 
 uncombined with the water. Solutions of sodium carbonate, 
 Na 2 COa, and potassium carbonate, E^COs, have an alkaline 
 
214 CARBON OXIDES, CYANIDES 
 
 reaction. Explain this as has been done for alkaline phos- 
 phates. 
 
 Solutions of carbon dioxide in water react with bases to 
 form carbonates. The effect of the gas on lime water, 
 Ca(OH) 2 , has been referred to several times (p. 11). 
 
 Carbon dioxide can be liquefied by pressure and the 
 liquid is sold in strong steel cylinders for use in soda-water 
 fountains and for carbonating beverages. The pressure 
 in these cylinders is from 60 to 70 atmospheres, or 900 to 
 1000 pounds to the square inch. If the liquid is allowed to 
 escape into a strong cloth bag a part of it evaporates at 
 once as the pressure is removed and the remainder is cooled 
 to so low a temperature that it freezes to a white, snow- 
 like solid. The temperature of this solid is 79, the boiling 
 point, or rather subliming point, of solid carbon dioxide. 
 In other words, the vapor pressure of solid carbon dioxide is 
 760 mm. at - 79. The melting point of the solid is - 56.4, 
 more than 20 above the boiling point. The vapor pres- 
 sure at the melting point is 5.1 atmospheres. 
 
 Solid carbon dioxide is a very convenient means of 
 securing low temperatures in the laboratory. 
 
 Carbon Bisulfide. When the vapor of sulfur is passed 
 over heated charcoal the sulfur combines with the carbon 
 to form carbon disulfide, CS 2 . The preparation is now 
 carried out in an electric furnace. 
 
 Carbon disulfide, when pure, is a colorless liquid which 
 boils at 47 and gives off a considerable amount of vapor 
 at ordinary temperatures. The kindling temperature is 
 very low, and mixtures of the vapor with air are highly 
 explosive. For this reason carbon disulfide must be handled 
 with extreme care. This inflammability interferes with 
 its use as a solvent for fats and for some purposes to which 
 it is otherwise well adapted. It is used in vulcanizing 
 india-rubber, in rubber cements and some times as a poison 
 to kill rats, ground squirrels and moths. 
 
CYANIDES 215 
 
 Potassium Ferrocyanide. When a mixture of slaughter 
 house refuse, rich in nitrogenous carbon compounds, is 
 heated with potassium carbonate, K 2 CO 3 , and iron turn- 
 ings the elements unite to form potassium ferrocyanide, 
 K 4 Fe(CN) 6 . This is a compound formed by the union of 
 potassium cyanide, KCN, and ferrous cyanide, Fe(CN) 2 , 
 but it is also to be considered as a salt of the acid, 
 H 4 Fe (CN) 6 , called hydroferrocyanic acid. Potassium ferro- 
 cyanide dissolves readily in water and crystallizes from 
 the solution as a yellow hydrate, having the composition, 
 K 4 FeC 6 N 6 .H 2 O. 
 
 Cyanides. When potassium ferrocyanide is heated with 
 sodium the iron is replaced by the sodium, giving a mixture 
 of potassium cyanide, KCN, and sodium cyanide, NaCN: 
 
 K 4 FeC 6 N 6 + 2Na = 4KCN + 2NaCN + Fe 
 
 A solution of these cyanides will dissolve metallic gold, 
 if used in the presence of air, and such a solution is exten- 
 sively used in extracting gold from its ores. 
 
 Hydrocyanic Acid or Prussic Acid. If a solution of 
 potassium ferrocyanide, K 4 FeC 6 N 6 or of potassium cyanide, 
 KCN, is mixed with dilute sulfuric acid and distilled, 
 hydrocyanic acid, which is volatile, will pass over. Pure 
 hydrocyanic acid is a volatile liquid which boils at 26.5. 
 It is one to the quickest, most powerful poisons known. 
 The cyanides are also very poisonous. 
 
 A dilute solution of hydrocyanic acid is sometimes used 
 in medicine. 
 
 Complex Cyanides. Many of the cyanides of heavy 
 metals are insoluble in water but most of these cyanides 
 will dissolve in a solution of potassium cyanide. In the 
 resulting solution the heavy metal enters into a complex 
 group which reacts as a whole and no longer shows the 
 characteristics of the heavy metal which it contains. Thus 
 potassium ferrocyanide, K 4 FeC 6 Ne, will give no precipitate, 
 
216 CARBON OXIDES, CYANIDES 
 
 with sodium hydroxide, NaOH, ammonium sulfide, (NH 4 ) 2 S, 
 or with other reagents which react readily with ferrous 
 sulfate, FeSC>4, or with ordinary ferrous salts. 
 
 If an electric current is passed through a solution of ferrous 
 sulfate, FeSO 4 , the ferrous ion, Fe ++ , is carried toward 
 the negative pole, or cathode, and the sulfate ion, SO 4 = , 
 is carried toward the positive pole, or anode. If the current 
 is passed through a solution of potassium ferrocyanide, 
 however, the iron is not carried toward the cathode but 
 it is carried, instead, with the cyanogen, CN, toward the 
 anode and only the potassium travels toward the cathode. 
 This is best explained by supposing that the ions in the 
 solution are not 4K + , Fe ++ and 6CN~, as might have been 
 expected, but that they are 4K+ and FeC 6 N 6 E . In other 
 words the iron and cyanogen unite to form a complex ion 
 which is called the ferrocyanide ion. 
 
 Many other complex cyanides are known. Among these 
 may be mentioned silver argenticyanide, KAgC 2 N 2 (or 
 KCN.AgCN), which is used in silver plating, and potas- 
 sium ferricyanide, K 3 FeC 6 N 6 , or 3KCN.Fe(CN) 3 , a red 
 salt prepared by oxidizing potassium ferrocyanide. The 
 last salt is used in preparing blue-print paper. 
 
 SUMMARY 
 
 Carbon monoxide is formed when carbon dioxide passes 
 over hot carbon. The reaction is endothermic. 
 
 Oxalic acid when heated with concentrated sulfuric 
 acid gives carbon monoxide and carbon dioxide. The 
 carbon dioxide may be absorbed by soda lime. 
 
 Carbon monoxide burns with a blue flame. It is very 
 poisonous. It is a constituent of water gas. 
 
 Carbon dioxide is formed by burning carbon. It is 
 prepared by the action of an acid on a carbonate. 
 
 Water dissolves about its own volume of carbon dioxide 
 
EXERCISES. CARBON OXIDES 217 
 
 whether under high or low pressure (Henry's Law) and so 
 dissolves a greater weight under high pressures. The 
 solution contains carbonic acid, which is a very weak acid. 
 The salts of the alkali metals have an alkaline reaction in 
 solution. 
 
 The subliming point of pure carbon dioxide is 79. 
 The melting point is more than 20 higher. Liquid carbon 
 dioxide is possible only under pressure. 
 
 Carbon disulfide is prepared by passing sulfur vapor over 
 hot charcoal. It is poisonous and very inflammable. It 
 is used in vulcanizing rubber, and sometimes as a poison. 
 
 Potassium ferrocyanide is prepared by heating nitro- 
 genous matter with potassium carbonate and iron. 
 
 A mixture of potassium and sodium cyanides is prepared 
 by heating potassium ferrocyanide with sodium. It is used 
 in extracting gold from its ores. 
 
 Hydrocyanic or prussic acid and the cyanides are very 
 poisonous. 
 
 Potassium cyanide combines with cyanides of the heavy 
 metals to form complex cyanides in which the heavy metal 
 forms a part of the anion. Solutions of such complex 
 cyanides often fail to give the ordinary reactions used to 
 detect the heavy metals which they contain. 
 
 Potassium argenticyanide is used in silver plating. 
 
 EXERCISES 
 
 1. Write the equation for the action of dilute sulfuric acid on 
 potassium ferrocyanide. 
 
 2. When a slightly diluted sulfuric acid is heated with potassium 
 ferricyanide, carbon monoxide is formed. The other products are 
 ammonium sulfate, potassium sulfate and ferric sulfate. Write 
 the equation. Notice that two molecules of the ferricyanide 
 must be used (why?) and enough water to furnish the oxygen of the 
 carbon monoxide. 
 
 3. In what proportion by volume should carbon monoxide 
 and oxygen be mixed for explosion? When cold what will be the 
 
218 CARBON OXIDES, CYANIDES 
 
 volume of the gaseous product as compared with volume of the 
 original gases? 
 
 4. If the mixture of carbon monoxide and oxygen gives a tem- 
 perature of 2500, what will be the volume of the carbon dioxide 
 formed at that temperature as compared with the volume of the 
 mixed gases at 20 before the explosion? 
 
 5. Sodium carbonate, Na 2 C0 3 , is hydrolyzed by water, in part, 
 giving hydrocarbonate, HC0 3 ~ and hydroxide, OH~, ions. Write 
 the equation for the reaction- between the ions of water and the 
 ions of sodium carbonate showing the ions which result. What 
 will be the reaction of the solution? Confirm this by testing a 
 solution of sodium carbonate with litmus paper. 
 
 6. What per cent of the heat energy of carbon is lost when it is 
 burned only to carbon monoxide? 
 
CHAPTER XX 
 
 CARBOHYDRATES, ALCOHOLS, ACIDS, BREAD, PRO- 
 TEINS, DIGESTION, ANTITOXINS, ALKALOIDS, DYES 
 
 Carbohydrates. Among the many thousands of com- 
 pounds which contain carbon, hydrogen and oxygen, those 
 of one of the most important groups are called carbohydrates 
 because the hydrogen and oxygen in them are in the same 
 proportion as in water. The name seems to imply that 
 they are hydrates of carbon, that is, that they are formed 
 by the union of carbon with water. Such a point of view 
 is not justified, for they are not formed in nature, or arti- 
 ficially, by the direct union of carbon with water, and when 
 heated, while carbon and water are formed by their decom- 
 position, many other substances are formed as well. 
 
 The carbohydrates include, especially, cellulose, starch 
 and many different sugars. 
 
 Cellulose, is represented by the formula (C 6 HioO 5 ) n . 
 No means has been discovered for determining the value of 
 n in this formula, because cellulose cannot be vaporized 
 without decomposition and there is no simple solvent in 
 which it might be dissolved and its molecular weight deter- 
 mined by the lowering of the freezing point or rise of the 
 boiling point of its solution as is done with many other 
 compounds which cannot be vaporized. 
 
 Cellulose forms the larger part of the woody fiber of trees 
 and all kinds of plants. 
 
 In the form of grass, clover, alfalfa and the hay or silage 
 made from these or from corn it is an important food for 
 herbivorous animals. It is a constituent of many foods 
 
 219 
 
220 CARBOHYDRATES, ALCOHOLS, ETC. 
 
 used by man but apparently it is not digested and utilized 
 to any appreciable extent. 
 
 Coal was formed largely from the woody fiber of plants 
 which grew ages ago. Coal and wood form, of course, our 
 most important fuels. There is some probability that 
 natural gas and petroleum were formed, in part, from 
 cellulose. 
 
 Cotton, linen, hemp and other vegetable fibers used in 
 the manufacture of cloth, ropes and twine are largely com- 
 posed of cellulose. 
 
 Paper is almost entirely cellulose. The cheaper grades, 
 used for printing the daily papers, are made from wood. 
 Better kinds of paper are made from linen rags and other 
 fibrous materials. The materials are bleached and purified 
 by the use of various chemicals and mixed with water to a 
 thin pulp which can be spread out in a uniform layer, which 
 is then pressed and dried. 
 
 Nitrocellulose; Gun Cotton; Lacquers; Collodion; Arti- 
 ficial Silk. When cotton, which is nearly pure cellulose, 
 is treated with a mixture of concentrated sulfuric and 
 nitric acids it is converted into a mixture of compounds 
 called usually nitrocellulose, but more correctly cellulose 
 nitrate. These are formed by the replacement of hydroxyl 
 groups, OH, by the nitrate group, NO 3 , just as sodium 
 nitrate is formed by the replacement of the hydroxyl group 
 of sodium hydroxide: 
 
 NaOH + HNO 3 = NaN0 3 + HOH 
 Ci 2 H 14 O 4 (OH) 6 + 6HNO 3 = C 12 H 14 O 4 (NO 3 ) 6 + 6HOH 
 
 Cellulose hexanitrate 
 
 Cellulose hexanitrate is the powerful explosive known as 
 gun-cotton and also used as the basis of the smokeless 
 powders. 
 
 Other nitrates are formed by the replacement of a smaller 
 number of hydroxyl groups. Solutions of some of these in 
 
CELLULOID. STARCH 
 
 221 
 
 amyl acetate or other solvents are excellent lacquers for 
 brass. Collodion is a solution in ether and alcohol. Arti- 
 ficial silk is made from cellulose nitrate or acetate but if the 
 former is used it is subjected to a treatment which removes 
 the nitrate group and makes it less inflammable. 
 
 Celluloid is a mixture of some of the cellulose nitrates 
 with camphor. It is highly inflammable but not explosive 
 in the ordinary sense. 
 
 $fr$ 
 
 o^x^.R OWQf* 
 
 FIG. 41. A, potato starch; B, rice starch; C, wheat starch 
 (X 160). After Allen. 
 
 Starch also has the formula (C 6 H 10 O 5 )n and its molecular 
 weight is unknown It is found in the form of granules of 
 various sizes and shapes in wheat, maize, rice, sago, tapioca 
 and many other substances used as articles of food. It is 
 the most important non-nitrogenous compound in articles 
 of human diet and it furnishes a considerable portion of the 
 heat of our bodies and of the muscular energy with which we 
 move and do work. 
 
'222 CARBOHYDRATES, ALCOHOLS, ETC. 
 
 The forms of the granules, some of which are illustrated 
 in Fig. 41 have no connection with the chemical composition, 
 which seems to be the same in all plants. A very thin, outer 
 shell covering the granules is probably of a. somewhat dif- 
 ferent character from the starch, but it has not been found 
 possible to separate and examine it. 
 
 When foods containing starch are cooked the cell walls 
 are burst and the starch forms a soft, pulpy mass, which is 
 easily attacked by the digestion fluids. Pure starch forms 
 a paste with hot water. With larger amounts of water it 
 gives a slightly opalescent, colloidal solution. 
 
 Starch gives an intense blue color with a solution of iodine 
 in potassium iodide. This is used as a sensitive test for 
 starch or for iodine. 
 
 Sucrose or Cane Sugar. Beet Sugar. The juices of 
 isugar cane, beets, the sap of maple trees and nearly all 
 fruits contain a crystalline, easily soluble compound hav- 
 ing the composition C^H^On. It is commonly known 
 merely as sugar, more accurately as sucrose, or cane 
 sugar; It is manufactured chiefly from sugar cane and 
 sugar beets. 
 
 The juice of the sugar cane is pressed out with powerful 
 rolls. It is then concentrated to the point of crystalliza- 
 tion by evaporation under diminished pressure. The reduc- 
 tion of the pressure causes the solution to boil at a lower 
 temperature and there is much less decomposition of the 
 sugar than if the solution were boiled down at atmospheric 
 pressure. The crystals of sugar which are deposited on 
 cooling the concentrated solution are separated from the 
 syrup (molasses) by means of a centrifugal strainer. 
 
 The sugar from beets, when properly purified, is identical 
 with cane sugar. 
 
 Maple sugar is allowed to retain substances which give 
 to it a desirable flavor. The pure sugar is the same as that 
 in sugar cane or sugar beets. 
 
SUGARS 223 
 
 Invert Sugar. If ordinary sugar is warmed for a short 
 time with a dilute acid it takes up water and is converted 
 into a mixture of equal parts of two other sugars, glucose 
 and fructose. The same change can be effected by certain 
 organic ferments and in other ways: 
 
 C 12 H 22 On + H 2 = C 6 H 12 6 + C 6 H 12 6 
 
 Sucrose Glucose Fructose 
 
 Invert sugar 
 
 Sucrose rotates the plane of a ray of polarized light to 
 the right. Glucose also rotates the plane of the ray to the 
 right but fructose rotates it to the left to a greater degree 
 at ordinary temperatures and for this reason the mixture 
 causes a left-handed rotation and it is often called invert 
 sugar. Invert sugar is found in honey, in some fruit 
 juices and in syrups made from sugar cane or sorghum. 
 Owing chiefly to the fructose which it contains, it is sweeter 
 than ordinary sugar. 
 
 Glucose. The formation of glucose by the hydrolysis of 
 cane sugar has just been mentioned. It may also be pre- 
 pared by the hydrolysis of starch by boiling it with dilute 
 sulfuric or hydrochloric acid. Large quantities of the sugar 
 are manufactured in this way and sold in cheap candies and 
 in syrups, especially in the syrup known as corn syrup. 
 The acid is, of course, neutralized or removed. 
 
 In the disease called diabetes, sugar or starch is' con- 
 verted in part into glucose and eliminated from the body 
 in that form. This has given rise to an impression that 
 glucose is not a safe article of diet. Ordinary sugar is 
 converted partly into glucose during digestion and there 
 seems to be no scientific ground for thinking glucose any 
 more harmful than sucrose. 
 
 Maltose. When starch, which has been boiled to rupture 
 the granules, is mixed with malt and warmed to 65-70 
 it is largely converted into a sugar called maltose, which has 
 
224 CARBOHYDRATES, ALCOHOLS, ETC. 
 
 the same composition as cane sugar, Ci 2 H 22 On, but differs 
 from it in some of its properties: 
 
 2(C 6 H 10 O5)n + nH 2 = nCi 2 H 22 On 
 
 Starch Maltose 
 
 Malt is prepared by moistening barley and allowing it to 
 germinate. As the barley sprouts an enzyme called diastase 
 is formed. This is soluble in water and a small quantity of 
 it will convert a large amount of starch into sugar. Dias- 
 tase is one of a considerable number of organic substances 
 called enzymes, which act as catalytic agents. 
 
 Dextrin. If starch is moistened with very dilute nitric 
 acid and heated for sometime at 120 it is converted into an 
 easily soluble substance called dextrin. This is used in 
 making mucilage and for the backs of postage stamps. 
 
 Pectose. Pectin. Jelly. Fruits of nearly all kinds, 
 especially when not fully ripe contain an insoluble sub- 
 stance called pectose. When the fruits are boiled with 
 water the pectose is decomposed and yields a soluble sub- 
 stance called pectin. Pectin forms a jelly with sugar, in a 
 slightly acid solution. In making jelly, from one-half to 
 three-fourths of a cupful of sugar is added for each cupful 
 of fruit juice. The fruit juice should contain from 0.5 
 to 0.7 per cent of acid, calculated as tartaric acid. " The 
 boiling should not be continued too long after separating 
 the juice from the fruit, as the pectin seems to be slowly 
 destroyed by heat. 
 
 Ethyl Alcohol. When liquids containing sugar, such as 
 the juice of grapes, apples and other fruits, or syrups ob- 
 tained in the manufacture of sugars, are exposed to the 
 air they almost always acquire spores of yeast. These 
 grow and cause the fermentation of the sugar with the for- 
 mation of alcohol and carbon dioxide. The fermentation 
 seems always to be preceded by the change of the sugar to 
 invert sugar (see cane sugar, above). 
 
ALCOHOL 
 
 225 
 
 The solution of maltose which is prepared by warming 
 cooked starch with malt may also be fermented by yeast : 
 
 C 6 H 12 O 6 = 2C 2 H 6 O + 2CO 2 
 
 Glucose or Alcohol 
 
 fructose 
 
 Ci 2 H 22 O n 
 
 Maltose 
 
 H 2 - 4C 2 H 6 + 4C0 
 
 In manufacturing alcohol, corn meal or potatoes are first 
 thoroughly cooked and the cooked material is mixed with 
 
 Dilute 
 Alcohol - 
 
 Steam Jacket- 
 
 * To Condenser 
 
 Alcohol- 
 Free Wafer 
 
 FIG. 42. 
 
 about 10 per cent of its weight of malt and enough water 
 so that the resulting solution will contain about 10 per 
 cent of sugar. After warming for a short time to bring 
 
 15 
 
226 CARBOHYDRATES, ALCOHOLS, ETC, 
 
 about the action of the diastase on the starch the solution 
 is cooled to about the temperature of the hand and yeast is 
 added and the mixture is allowed to ferment for three or 
 four days. The fermented liquid is then distilled to sepa- 
 rate the alcohol from the large quantity of water present. 
 Alcohol boils at 78 and water at 100 and when a mixture 
 of the two is distilled the portions passing over first will 
 contain more alcohol than the original liquid. With a 
 " column" still constructed on the principle of the diagram, 
 (Fig. 42), by introducing the dilute alcohol near the center 
 the stronger and stronger alcohol distils upward from one 
 shelf to another while the water containing less and less 
 alcohol runs downward and finally leaves the still practi- 
 cally free from alcohol at the bottom. The alcohol which 
 reaches the top of the still may contain 90 per cent or 
 more of alcohol and less than 10 per cent of water. 
 
 Alcohol is used for burning, as a solvent in making 
 varnishes and in preparing pharmaceutical extracts and 
 tinctures. 
 
 "Denatured alcohol" contains substances which have 
 been added to render it unsuitable for use as a beverage. It 
 is sold free of tax and may be used for burning and in making 
 varnishes but it is much more poisonous than pure alcohol 
 and is unfit for drinking or for any medicinal use. 
 
 Acetic Acid, HC 2 H 3 O 2 . When dilute alcohol, such as 
 is formed by the fermentation of cider, wine or other sac- 
 charine liquids, is exposed to the air in loosely closed casks 
 it almost invariably acquires bacteria from the air, which 
 causes a second fermentation to acetic acid. The com- 
 mercial product is called vinegar. Alcoholic fermentation, 
 which is caused by yeast, takes place in closed vessels 
 from which the air -is excluded. The acetic fermentation 
 requires the presence of air to furnish the oxygen required 
 for the oxidation of the alcohol: 
 
 C 2 H 5 OH + O 2 = C 2 H 3 O.OH(or HC 2 H 3 O 2 ) + H 2 O 
 
FATS, SOAP 227 
 
 In vinegar factories beech-wood shavings are inoculated 
 with the bacteria which cause -the acetic fermentation and 
 the dilute alcohol is permitted to run slowly over these in 
 such a manner as to expose a large surface of the liquid to 
 the combined action of air and the bacteria. Good vinegar 
 should contain 4 per cent of acetic acid. 
 
 Fats. Such substances of lard, tallow, olive oil, cotton 
 seed oil, and butter are composed almost entirely of 
 compounds called fats. All of the fats contains acids 
 called fatty acids, whose hydrogen has been replaced 
 by the group C 3 H r ,, called glyceryl. This group is charac- 
 teristic of glycerol (commonly called glycerine) C 3 H 5 - 
 (OH) 3 . The three most common compounds found in 
 fats are derived from the three fatty, acids, stearic 
 acid, HCi8H 35 O2, palmitic acid, HCi 6 H 3 iO2, and oleic 
 acid, HCisHssC^. The compounds derived from these are 
 stearin, C 3 H 5 (Ci8H 35 O2)3, palmitin, C 3 H 5 (Ci 6 H 3 iO 2 ) 3 and 
 olein, C 3 H 5 (Ci8H33O 2 )3. 
 
 Soap. When a fat is heated with a concentrated solution 
 of sodium hydroxide, NaOH, it is decomposed, forming 
 the sodium salt of the acid and glycerol, C 3 H 5 (OH) 3 : 
 
 2 ) 3 + 3NaOH = 3NaCi 8 H 35 O 2 + C 3 H 5 (OH) 3 
 
 Stearin Sodium stearate 
 
 This process is called saponification and the salts formed 
 are called soaps. Soaps dissolve more or less readily in 
 water and give solutions which will emulsify fats and oils. 
 if olive oil or kerosene is shaken with water and the mixture 
 is allowed to stand for a short time the oil separates as 
 a layer floating on the water. But if a solution of soap 
 is added and the mixture shaken again an emulsion is formed 
 from which the oil and water will separate very slowly in- 
 deed or not at all. On rubbing soiled clothes with soap 
 and water the cleansing of the cloth depends partly on the 
 formation of an emulsion with greasy matters on the 
 
228 CARBOHYDRATES, ALCOHOLS, ETC. 
 
 fabric, because the soapy solution can wet such substances 
 and emulsify them, while pure water cannot. The emulsion 
 and particles of dirt, which are also wet by the solution, 
 can then be rinsed away with the water. 
 
 Glycerol or Glycerine, C 3 H 5 (OH)3, is obtained as a 
 by-product in making soap from fats. It is used for a 
 variety of purposes but chiefly in the manufacture of nitro- 
 glycerine. 
 
 Nitroglycerine is made by treating glycerine with a 
 mixture of nitric and sulfuric acids. 
 
 C 3 H 5 (OH) 3 + 3HN0 3 = C 3 H 5 (N0 3 ) 3 + 3H 2 O 
 
 Glycerine Nitroglycerine 
 
 Nitroglycerine is mixed with sawdust or some other 
 porous substance to make dynamite. Either nitroglycerine 
 or dynamite is exploded by a detonating cap of fulminate of 
 mercury. The explosion is due to the combination of the 
 oxygen which it contains with its carbon and hydrogen 
 forming a large volume of carbon dioxide and steam from a 
 small volume of the liquid. The combination is attended 
 with considerable evolution of heat and this causes the 
 expansion of the gases and intensifies the explosion. Some 
 nitric oxide, NO, is formed by the explosion and confined 
 spaces where nitroglycerine or dynamite is exploded require 
 ventilation before entering them after the explosion. 
 
 Phenol or Carbolic Acid, C 6 H 5 OH, is obtained from 
 coal tar and is also prepared synthetically. It is used as a 
 germicide and disinfectant but is effective only when applied 
 directly. The vapor is not concentrated enough to be of 
 value. Because of its strong disagreeable odor, uninformed 
 persons are often given a misleading sense of security by 
 its use. It is the chief active constituent in the " coal-tar 
 dips" used in the care of sheep. 
 
 Tartaric Acid, I&C^Oe When grape juice is allowed 
 to ferment in the manufacture of wine an acid salt of tar- 
 
BREAD 229 
 
 taric acid which is known as cream of tartar, KHC 4 H 4 6 , 
 separates because the salt is much less soluble in dilute 
 alcohol than it is in water. The chemical name of the salt 
 is acid potassium tartrate. It has a sour taste and reacts 
 readily with bases or with carbonates, giving neutral salts: 
 
 KHC 4 H 4 O 6 + NaOH = KNaC 4 H 4 O 6 + H 2 O 
 
 Sodium potassium tartrate 
 
 KHC 4 H 4 O 6 + NaHCO 3 = KNaC 4 H 4 O 6 + CO 2 + H 2 O 
 
 Bread. The manufacture of a palatable food from wheat 
 flour or from flour prepared from other cereals depends 
 largely upon securing a fine cellular structure of the cooked 
 material. Such a structure permits the easy access of the 
 saliva and other digestive fluids to all parts of the substance . 
 and in this way promotes its digestion. In making bread 
 the flour is mixed with water or milk, or both, and with some 
 yeast, to a stiff dough and the whole is thoroughly kneaded 
 by hand or by some mechanical device to secure the fine 
 cellular structure which is required. The yeast acts on 
 the small amount of sugar present fermenting it to alcohol 
 and carbon dioxide. The latter distends the little cells or 
 interstices in the dough, causing it to "rise." After a time 
 the kneading is repeated to secure a more uniform structure 
 and after being allowed to rise once more the dough is 
 baked in an oven. It is important that the materials used 
 shall be slightly warm and that the bread shall be kept at a 
 temperature of 21-26 (70-80 F.) during the fermenta- 
 tion or rising. At a lower temperature the yeast acts too 
 slowly and at higher temperatures it may be killed. 
 
 The alcohol formed by the fermentation partly escapes 
 during the baking of the bread but the larger portion of 
 it is retained. The amount formed is, of course, small. 
 
 Baking Powders. If cream of tartar (acid potassium 
 tartrate, KHC 4 H 4 O 6 ) and baking soda (acid sodium car- 
 bonate, NaHCO 3 ) are mixed dry no reaction occurs, but 
 
230 CARBOHYDRATES, ALCOHOLS, ETC. 
 
 on adding water they react to form potassium sodium 
 tartrate, carbon dioxide and water (see tartaric acid, above). 
 In a similar manner alum (potassium aluminium sulfate, 
 KM (80)4) 2), will not react with baking soda when dry, but on 
 the addition of water carbon dioxide is liberated in accord- 
 ance with the equation: 
 
 KAl(S0 4 )-> + 3NaHCO 3 = KNaSO 4 + Na 2 S0 4 +3CO 2 - 
 
 Al(OH) 
 
 Proteins. Albumin. Casein. Gluten. Complex com- 
 pounds containing carbon, hydrogen, oxygen, nitrogen and 
 usually sulfur or phosphorus, called proteins, are found 
 in all living plants and animals. The most familiar of 
 these are albumin, which forms the larger part of the white 
 of an egg, and casein, the chief constituent of the curd which 
 can be separated from skim milk. Compounds of similar 
 composition are found in the muscular fiber of meat and in 
 the gluten of wheat flour, which remains when the flour is 
 kneaded between the fingers in a stream of running water, 
 to wash away the starch. 
 
 Digestion. Formation of Tissues. Production of Heat 
 and Energy. Food which is eaten performs in the body two 
 or three distinct functions. A part of the food is oxidized 
 to carbon dioxide and water, a process which corresponds 
 closely to the burning of a fire, and the heat generated 
 maintains the temperature of the body above the tempera- 
 ture of the air which surrounds it. A part of the energy of 
 the food is also converted into the muscular energy with 
 which we move and do work. If these were the only func- 
 tions to be served in the body our food might consist 
 exclusively of compounds of carbon, hydrogen and oxygen, 
 that is, of such compounds as the fats and carbohydrates. 
 
 1 Notice that enough sodium is required to replace the aluminium, 
 which is trivalent, and that the aluminium forms the hydroxide, Al(OH)s. 
 The rest of the equation follows from these facts. 
 
GROWTH OF PLANTS 231 
 
 It will be seen from this, too, why foods containing fats are 
 more suitable in cold than in warm climates. 
 
 The tissues of the body are also constantly broken down 
 and must be restored and during growth new tissues must' 
 be produced. Since the proteins form the most important 
 part of the tissues, it is evident that foods must always 
 contain nitrogen, sulfur and phosphorus as well as carbon, 
 hydrogen and oxygen. All of these elements must also be 
 present in the food in such a form that they can be digested 
 and assimilated. 
 
 In the process of digestion proteins are decomposed into 
 simpler compounds which are soluble and which can then 
 be transported by the blood to parts of the body where they 
 are needed for building or restoring tissues. Fats are 
 emulsified and brought into the circulation in that form. 
 By respiration the oxygen of the air is brought to one side 
 of the thin membranes of the lungs while the venous blood 
 is brought to the other side of the membranes by circulation 
 from the heart. The blood brings with it carbon dioxide 
 formed by the oxidation of tissues and compounds in the 
 body. Through the membranes of the lungs it gives up 
 this carbon dioxide and absorbs oxygen in its place. The 
 oxygen is carried by the blood through the arterial circula- 
 tion all over the body and is used in the oxidation which 
 produces heat and muscular energy. 
 
 Growth of Plants. From what has been given in the last 
 paragraph it can be seen that the animal body secures 
 the energy to maintain its existence by the oxidation of 
 food. The energy for the growth of plants is obtained by a 
 radically different process. The carbon for the plant is 
 taken directly from the carbon dioxide of the air. The 
 energy to separate the oxygen from the carbon of the carbon 
 dioxide is absorbed from the sunlight by the leaves of the 
 plant. The nitrogen for the plant must be furnished by the 
 soil, either directly in the form of nitrates, or ammonia, or 
 
232 CARBOHYDRATES, ALCOHOLS, ETC. 
 
 through the agency of nitrogen-fixing bacteria which thrive 
 in the roots of some leguminous plants, such as clover and 
 alfalfa. Potassium, phosphorus, sulfur and other mineral 
 constituents necessary for the growth of plants must also be 
 supplied by the soil. More than ninety per cent of the 
 weight of growing plants is derived from the water of the 
 soil and the carbon dioxide of the air. 
 
 Toxins and Antitoxins. It is well known that such dis- 
 eases as diphtheria, typhoid fever, tuberculosis, yellow fever, 
 malarial fever and many others are caused by minute 
 organisms called bacteria. Some of these organisms when 
 they find a lodgment and grow in the body produce poisons 
 called toxins. These are often very virulent and may pro- 
 duce death. It has been discovered, however, that under 
 the stimulus of the toxin the body produces an antidote 
 called an antitoxin. Very little is known about the exact 
 nature or composition either of the toxins or antitoxins, 
 but it has been discovered that in some cases the antitoxin 
 may be developed in animals and used as an antidote for 
 toxins in the human body. Thus by inoculating a horse 
 with the bacteria which cause diphtheria the antitoxin for 
 the disease may form in large quantities in the blood of 
 the horse, and the serum from the blood, if injected into a 
 person suffering from the disease will, in most cases, effect 
 a cure. 
 
 Alkaloids. A number of plants produce basic substances 
 containing carbon, hydrogen, nitrogen and usually oxygen, 
 which are called alkaloids. This name is given to them 
 because they combine with acids to form salts, as alkalies 
 and other bases do. Some of the alkaloids are powerful 
 poisons and most of them produce marked physiological 
 effects. The best known are strychnine, morphine, nico- 
 tine, cocaine, atropine and quinine. Nearly or quite all 
 alkaloids are bitter; some of them, especially strychnine 
 and quinine, intensely so. 
 
ALKALOIDS, DYES 233 
 
 Strychnine is a very powerful poison but it is also used 
 in small doses as a heart stimulant. 
 
 Morphine is obtained from the poppy and is used to pro- 
 duce sleep. It is the chief constituent of opium and 
 laudanum. Morphine and opium are among the most 
 dangerous of the habit-forming drugs. 
 
 Nicotine is a volatile, liquid alkaloid found in tobacco. 
 It is very poisonous, but comparatively small quantities of 
 it are volatilized in smoking. 
 
 Cocaine is used to produce a local anesthesia in minor 
 surgical operations. 
 
 Atropine has been used by oculists to cause a widening 
 of the pupil of the eye so that the retina may be examined to 
 better advantage. It is now largely displaced by other 
 compounds which are more suitable. 
 
 Quinine is obtained from Peruvian bark. It is specific 
 for malarial fever and is sometimes used for other pur- 
 poses in medicine. 
 
 Dyes. From early times, natural substances obtained 
 from animal and vegetable sources have been used to give 
 beautiful colors to skins and cloths. The most common 
 natural dyes of this sort are cochineal, Turkey red, indigo, 
 fustic and logwood. In 1856 a young English chemist, 
 William H. Per kin, discovered that a beautiful dye, called 
 mauve, could be made from aniline, a compound which 
 can be prepared from the benzene of coal tar. Not long 
 afterward alizarin, the compound which gives the red color 
 of Turkey red, was prepared artificially from another 
 compound found in coal tar and it was very soon found 
 that the artificial alizarin could be made more cheaply 
 than the natural product. More recently indigo has been 
 made artificially at a profit and since 1856 about a thousand 
 dyes which are made artificially have been patented. Most 
 of these are different from any of the natural dyes and 
 almost every possible shade of color has been produced. 
 
234 CARBOHYDRATES, ALCOHOLS, ETC. 
 
 Because the first of the artificial dyes was made from 
 aniline, the artificial dyes have often been called " aniline 
 dyes," a name which is not very correct. The designation 
 "coal-tar dyes" is more proper, as nearly all of the artificial 
 dyes are made from compounds found in coal tar. Many 
 of the dyes at first discovered fade rapidly on exposure to 
 the light and some of them dissolve or "run" when the 
 cloth is washed. This has given a popular impression that 
 the artificial dyes are inferior to those from natural sources, 
 but this is by no means true of all of them, and such dyes 
 as alizarin and indigo are exactly the same when they are 
 made artificially as when obtained from madder root or 
 the indigo plant. 
 
 SUMMARY 
 
 Carbohydrates are compounds of carbon containing 
 hydrogen and oxygen in the same proportions as in water. 
 
 Cellulose is the principal constituent of wood, of cotton, 
 linen, hemp and other fibers, and of paper. 
 
 Nitrocellulose is used in gun cotton, smokeless powder, 
 lacquers, collodion and celluloid. 
 
 Starch is found in all cereals and is an important con- 
 stituent of foods prepared from these. 
 
 Sucrose is made from sugar cane, sugar beets, and maple 
 syrup. It is changed to invert sugar, a mixture of glucose 
 and fructose, by dilute acids. 
 
 Glucose is made by the action of acids on starch. It 
 is the principal constituent of corn syrup. 
 
 Maltose is made by the action of the diastase of malt on 
 starch. 
 
 Dextrin is made by warming starch moistened with a 
 little dilute nitric acid. 
 
 Pectose, which yields pectin on boiling with water, is 
 found in most fruits. It forms a jelly with acids and sugar. 
 
SUMMARY. CARBOHYDRATES, ETC. 235 
 
 Ethyl alcohol is made by the action of yeast on liquids 
 containing maltose, glucose or cane sugar. The starch 
 of corn or potatoes is first changed to maltose by the use 
 of malt. 
 
 Acetic acid and vinegar are prepared by the oxidation 
 of alcohol under the influence of bacteria. 
 
 Fats contain stearin, palmitin, olein and other compounds 
 of organic acids with glyceryl, the radical of glycerol. 
 They are saponified by alkalies, giving soap and glycerol. 
 
 Soap helps water to emulsify greasy substances so 
 that they can be removed. 
 
 Nitroglycerine is glyceryl nitrate. Dynamite is a mixture 
 of nitroglycerine with infusorial earth or some other ab- 
 sorbent material. 
 
 Phenol or carbolic acid is found in coal tar and it is also 
 manufactured. It is used as a germicide. 
 
 Tartaric acid is made from cream of tartar, which sepa- 
 rates from wines during fermentation. 
 
 Bread is raised by yeast, which causes the fermentation 
 of sugar to alcohol and carbon dioxide. 
 
 Baking powders are dry mixtures of acid sodium car- 
 bonate with cream of tartar, alum, or some other compound 
 which will liberate carbon dioxide from the baking soda 
 when the mixture is moistened. 
 
 Albumin, casein, gluten and other proteins are essential 
 constituents of foods and are required for restoring tissues 
 and for growth. 
 
 In digestion foods are dissolved and partly decomposed 
 to prepare them for introduction into the blood and for 
 use in restoring the tissues and maintaining the warmth 
 and muscular energy of the body. 
 
 Plants utilize the carbon dioxide of the air and the energy 
 of the sun and store energy. Animals dissipate the energy 
 of the food which they eat. 
 
 Toxins are poisons developed in the progress of disease. 
 
236 CARBOHYDRATES, ALCOHOL, ETC. 
 
 Antitoxins are antidotes for toxins which are spontaneously 
 developed in men or animals affected by a disease. 
 
 Alkaloids are basic compounds found in plants. The 
 most common are strychnine, morphine, nicotine, cocaine, 
 atropine and quinine. 
 
 Dyes were formerly mostly of vegetable origin, but the 
 larger part of them are now made in chemical factories. 
 
 EXERCISES 
 
 1. How many pounds of glucose could be made from a bushel 
 of corn containing 60 per cent of starch? A bushel of corn weighs 
 56 pounds 
 
 2. Assuming that 90 per cent alcohol has a specific gravity of 
 0.79 and that a gallon of water weighs 8.3 pounds, how many 
 gallons of 90 per cent alcohol could be made from a bushel of 
 corn? How many gallons of 4 per cent acetic acid? 
 
 3. Assuming that a grease used for soap consists of palmitin, 
 how much sodium hydroxide will be required to give 100 grams of 
 soap? , 
 
 4. In what proportions should baking soda and cream of tartar 
 be mixed in a baking powder? If tartaric acid were used in 
 place of cream of tartar, what would be the proportions to use? 
 
 5. What weight of alum, KA1(S0 4 ) 2 .12H 2 0, should be mixed 
 with 100 grams of baking soda in a baking powder? 
 
CHAPTER XXI 
 GROUP IV: SILICON, TIN AND LEAD 
 
 The Carbon Family of Elements. Carbon seems to be 
 the element on which the properties of the living matter of 
 plants and animals chiefly depend. The second element of 
 the same family is silicon and this is equally important as 
 the most abundant element after oxygen, in the solid crust 
 of the earth. Silica, the dioxide of silicon, SiO 2 , and sili- 
 cates, salts of acids of which silica is the anhydride, form 
 a very large proportion of the soil and of the minerals and 
 rocks which are most abundant. It is estimated that the 
 element silicon forms one-fourth of that portion of the earth 
 which we have been able to examine. Two other elements 
 of the same family, tin and lead, are very useful common 
 metals. 
 
 Just as the trioxides, N 2 3 , P20 3 , As 2 O 3 , Sb 2 3 and Bi 2 3 , 
 are characteristic of the elements of the nitrogen family, 
 which are trivalent, the dioxides, CO 2 , SiO 2 , PbO 2 and 
 SnO 2 are characteristic of the quadrivalent elements of the 
 carbon family. 
 
 The relations among the atomic weights may be recalled 
 by the following table: 
 
 C 12 
 
 N 14 
 
 O 16 
 
 F 19 
 
 Si 28 
 
 P 31 
 
 S 32 
 
 Cl 35.5 
 
 Ge 72 
 
 As 75 
 
 Se 78 
 
 Br 80 
 
 Sn 118 
 
 Sb 120 
 
 Te 127.6 
 
 I 127 
 
 Pb 207 
 
 Bi 208 
 
 
 
 
 
 While arsenic, antimony and bismuth of the nitrogen 
 family are brittle and in some of their properties are still 
 
 237 
 
238 SILICON, TIN AND LEAD 
 
 closely related to the non-metals, tin and lead are malleable 
 and fairly good conductors of heat and electricity. Anti- 
 mony and bismuth are sometimes called half-metals, but tin 
 and lead are always classed as metals. 
 
 Occurrence of Silicon. Preparation. There is some 
 reason for believing that all of the carbon on the earth has 
 been at one time combined with oxygen and that we find 
 carbon in the free state in coal, graphite and the diamond 
 because of the power which plants have had to reduce the 
 carbon dioxide of the atmosphere and form compounds 
 which have since decomposed with the separation of free 
 carbon. Silicon dioxide is not reduced in the growth of 
 plants and in spite of the abundance of the element there 
 seems to be no process occurring in nature by which silicon 
 is liberated in the free state. It is found exclusively as 
 silicon dioxide, SiO2, and in silicates formed by the union 
 of silicon dioxide with other elements. 
 
 Silicon is prepared commercially by heating silicon dioxide 
 with carbon in an electric furnace : 
 
 SiO 2 + 2C = Si + 2CO 
 
 Commercial silicon is gray and crystalline. It is used in 
 the manufacture of ferrosilicon, a very hard form of iron, 
 which is almost insoluble in acids. 
 
 Carborundum, SiC, is also prepared by heating a mixture 
 of silicon dioxide and carbon in an electric furnace. It 
 is a crystalline compound very much harder than emery and 
 has largely replaced the latter for grinding and cutting 
 purposes. 
 
 Silicon dioxide is found abundantly in nature in a 
 great variety of forms. The clear, transparent forms are 
 called rock crystal. Other forms, such as jasper, amethyst, 
 agate, rose quartz, smoky quartz and the like are colored by 
 minute amounts of other substances. The mineralogical 
 name for all common forms is quartz. Opal contains a 
 
GLASS 239 
 
 little water. Sand usually contains a large proportion of 
 quartz, because it is harder and less acted upon by water 
 than the other minerals in the rocks from which the sands 
 have been formed. 
 
 Quartz may be fused at a very high temperature with 
 the oxyhydrogen flame or in an electric furnace and can 
 be made into tubes, dishes and other forms of apparatus 
 which are useful in the laboratory. Such apparatus is 
 much less soluble than glass in water and also changes so 
 very little in volume when heated that it may even be 
 heated red hot and quenched in water without cracking. 
 
 Silicates. Nearly all of the very common rocks are 
 composed of mixtures of minerals which are silicates. 
 Many of them also contain quartz. The silicates are 
 derived from a series of hypothetical silicic acids none of 
 which are certainly known as definite compounds. They 
 all have the same anhydride, silicon dioxide. Granite is a 
 mixture of quartz, mica and feldspar. Other common 
 silicates are kaolin, the base of clay, garnet, talc or soapstone, 
 asbestos and meerschaum. 
 
 Glass. When silicon dioxide is heated with the oxides 
 or carbonates of such metals as potassium, sodium, calcium, 
 lead and some others, complex silicates are formed and if 
 the proportions are properly chosen the mixture solidifies 
 on cooling to a clear, transparent glass. The glass made in 
 this way is a mixture of silicates and has no definite melting 
 point. When heated it softens to a viscous liquid which 
 can be easily blown into large bulbs or cast or molded into 
 various forms. As it does not crystallize on cooling, but 
 remains clear, it is suitable for the great variety of uses 
 familiar to everyone. Window glass and the glass used for 
 bottles and other common articles is a silicate of sodium and 
 calcium with small quantities of other metals which are 
 present as impurities in the sand, lime and sodium carbon- 
 ate used in the manufacture. Flint glass contains lead 
 
240 SILICON, TIN AND LEAD 
 
 in place of calcium. It melts at a lower temperature, is 
 softer and has a higher index of refraction. Many different 
 kinds of glass are made for special purposes such as gage 
 glasses for steam boilers, beakers and flasks for laboratory 
 use, thermometers, lenses and imitations of diamonds and 
 other precious stones. 
 
 Soluble glass is a silicate of sodium prepared by fusing 
 sand and sodium carbonate: 
 
 Na 2 CO 3 + SiO 2 = Na 2 SiO 3 + CO 2 
 
 Sodium carbonate Sodium silicate 
 
 All kinds of glass dissolve in water to a slight degree, 
 though the ordinary forms are practically insoluble. 
 Sodium silicate, however, dissolves to a large extent, 
 giving a viscous solution which is used to cement glass 
 and porcelain, to fireproof cotton goods and to preserve 
 eggs. It is decomposed by acids, giving silicic acid which 
 may remain in solution as a colloid or may be precipitated 
 in a gelatinous form according to the way in which the 
 acid is added. 
 
 Colloidal Solutions. Many substances which are usually 
 insoluble may be obtained suspended in water in such fine 
 particles that they do not settle out as precipitates.- This 
 seems to be partly because the particles are so very fine 
 (from six to sixty millionths (0.000006 to 0.00006 mm.) 
 of a millimeter in diameter) and partly because they are 
 each composed of a number of molecules of the insoluble 
 substance gathered about some positive or negative ion. 
 The electrical charges of the ions cause the particles to 
 repel ;ih other and prevent them from uniting to form 
 large articles which would fall through the solution as a 
 precipitate. Solutions of this character are called colloidal 
 solutions. If the solution of colloidal silicic acid is placed 
 in the parchment sack of the apparatus shown in Fig. 43 
 and water is allowed to flow slowly through the bottle 
 around the sack, the salt, NaCl, and hydrochloric acid, 
 
COLLOIDS. DIALYSIS 
 
 241 
 
 FIG. 43. 
 
 HOI, of the solution will diffuse through the parchment and 
 
 be carried away while the silicic acid will remain behind. 
 
 In this way a nearly pure solution of colloidal silicic acid 
 
 may be obtained. Such a process is called dialysis and is 
 
 often used to separate colloids from substances which 
 
 dissolve in the molecular or ionic condition. The latter are 
 
 sometimes called crystalloids. The distinction seems to be 
 
 due to the fact that the 
 
 colloids consist of much larger 
 
 particles and for that reason 
 
 diffuse much more slowly and 
 
 may also be unable to pass 
 
 through the very fine pores 
 
 of the parchment. During 
 
 the process of digestion food 
 
 passes partly into a colloidal 
 
 solution and the colloids of 
 
 such solutions are separated 
 
 from simple molecalar compounds, which are also formed, 
 
 by the membranes of the digestive tract. 
 
 The very fine particles of clay form colloidal solutions 
 which render the clay plastic and make it possible to mold 
 the clay into the forms desired for the manufacture of 
 brick, tile, earthenware and china. Many other illustra- 
 tions might be given of the importance of colloidal solutions 
 for natural and industrial processes. 
 
 Occurrence of Tin. Tin is found hi nature as the min- 
 eral cassiterite, the dioxide, SnO 2 . Before the nineteenth 
 century the world's supply came almost exclusively from 
 Cornwall, England, where the ores have been mined ever 
 since the old Roman times. Tin is now obtained from 
 Banca, the East Indian islands and Tasmania. Only 
 small quantities have been found in the United States. 
 Metallic tin is readily obtained from cassiterite by heating 
 it with charcoal or coke, 
 
 16 
 
242 SILICON, TIN AND LEAD 
 
 Tin is a soft, white metal which melts at a low tem- 
 perature. It is less affected than any other common metal 
 by water, or by the combined action of water and air, 
 and for this reason it is used as a protective coating for 
 sheet iron for the manufacture of tin ware. 
 
 Solution Pressure of Metals. If a piece of a metal is 
 dipped in water or a dilute acid or alkali, atoms of the metal 
 tend to pass into the solution as positive ions. Thus 
 iron will give ferrous ions, Fe ++ , tin will give stannous ions, 
 Sn++, zinc will give zinc ions, Zn++. As these positive 
 ions pass into the solution, the piece of metal becomes nega- 
 tive because of the electrons (atoms of negative electricity) 
 given up by the atoms as they change to positive ions, 
 while the solution becomes positive. The amount of the 
 difference of potential between different metals and the 
 same solution or between the same metal and different 
 solutions varies with the character of the metal and that 
 of the solution. Unless there is some way provided for 
 the electrons to escape from the metal the escape of posi- 
 tive ions from its surface is almost instantly stopped by 
 this difference in potential. If two metals, such as iron 
 and tin, are placed in water or a dilute acid, the difference 
 in potential between the metal and the solution will be 
 greater for the iron than for the tin; in other words, the 
 iron will acquire a greater negative charge than the tin. 
 If the two metals are connected with a wire, the excess 
 of electrons in the piece of iron will pass into the tin. 
 New atoms of iron may then pass into solution till the 
 former difference of potential is restored. If there are 
 tin ions in the solution, electrons will escape from the tin 
 to these and discharge them, leaving the tin as metallic tin 
 on the surface. Under these conditions the stream of 
 electrons will be continued through the wire from the iron 
 to the tin and positive ions will make their way through 
 the solution toward the tin. In this manner an electric 
 
ELECTROMOTIVE SERIES 243 
 
 current can be maintained through the wire, the force 
 which drives the current coming from the fact that the 
 solution pressure of the iron, that is, the tendency of the 
 iron to pass into solution, is greater than the solution 
 pressure of the tin. 
 
 The characteristics of tin and iron which have just been 
 described have an important application in explaining 
 the conduct of the tinware used in the household. Pure 
 tin is scarcely affected by water or by the materials usually 
 employed in cooking. Iron, on the other hand, rusts slowly 
 in contact with water and air. A piece of iron in contact 
 with a piece of tin will rust more rapidly than the iron 
 alone because of the continual escape of electrons from the 
 iron to the tin. A new piece of tinware in which the coating 
 of tin over the iron is perfect does not rust and is not affected 
 by the ordinary liquids used in cooking. The moment a 
 little of the surface of the iron is exposed, however, the iron 
 rusts more rapidly than if the tin were not there. This 
 is in very marked contrast with the conduct of iron covered 
 with zinc (p. 283). 
 
 Electromotive Series. It has been pointed out that when 
 iron and tin are put in a solution and connected by means 
 of a wire a stream of electrons flows from the iron to the 
 tin while atoms of irqn pass into solution as ferrous ions, 
 Fe ++ . Because the iron has a stronger tendency than tin 
 to give up positive ions it is said to be more electro- 
 positive than tin. Zinc is still more electropositive than 
 iron, and copper is less electropositive than tin. In the 
 following table the common metals are arranged in the order 
 of the potentials which they assume toward each other. 
 Such a list is called the electromotive series: 
 
 Electropositive end: 
 Potassium 
 Sodium 
 Calcium 
 Magnesium 
 
244 
 
 SILICON-,. TIN, AND/ LEAD 
 
 Aluminium 
 Manganese 
 Zinc 
 Iron 
 Nickel 
 Lead 
 Tin 
 
 Hydrogen. 
 Bismuth 
 Copper 
 Antimony 
 Mercury 
 Silver 
 Platinum 
 Gold. 
 Electronegative end. 
 
 Direction ofCvrre 
 
 m\ 
 
 - 
 
 Zn 
 Electro 
 
 Positive 
 
 FIG. 44. 
 
 The position of a metal in the electromotive series is fixed 
 by determining the difference in potential between the 
 metal and a normal solution of the ions of the metal. 
 
 The relative positions of two metals in the series may be 
 determined by conflteGting stripe of the metals with a wire 
 
ALLOYS OF TIN 245 
 
 and determining the direction in which the current flows 
 (Fig. 44) through the wire when the strips of metal are 
 dipped in dilute acid. The stream of electrons (negative 
 atoms of electricity) flows from the electropositive metal 
 through the wire to the electronegative metal, i.e., in such 
 an arrangement the electropositive metal is negative and 
 the electronegative metal is positive. Thus if strips of iron 
 and tin are used the iron is electropositive and the positive 
 current of electricity flows toward it. 1 
 
 Alloys of Tin. Besides its use in the manufacture of tin- 
 ware tin is used in a number of alloys. Solder is an alloy 
 of tin and lead and a similar alloy is used for the tin-plate 
 for roofing purposes because lead is much cheaper than tin. 
 Babbitt metal is mostly tin, lead and antimony and is used 
 in the bearings of machinery because of its antifriction 
 properties. Many varieties of bronze contain tin and cop- 
 per as the most important ingredients. Tin is used in the 
 coils in which steam is condensed in preparing distilled 
 water because it is so little acted on by water. 
 
 Stannous Chloride, SnCl 2 .2H 2 O. Tin dissolves readily 
 in concentrated hydrochloric acid with the evolution of 
 hydrogen gas. The metal is bivalent in the chloride which 
 is formed but it has a strong tendency to take up oxygen 
 or more chlorine and pass over into a stannic compound. 
 This makes stannous chloride a good reducing agent and it 
 is often used for that purpose. 
 
 Stannic Chloride, SnCl 4 . Tin combines directly with 
 dry chlorine to form stannic chloride, in which the tin is 
 
 1 Confusion sometimes arises in the minds of students and others be^. 
 cause in the conventional system used by all physicists the current of 
 electricity is represented as flowing from the positive pole of a battery or 
 dynamo to the negative while the stream of electrons flows in the opposite 
 direction. When the older theories of electricity were developed there 
 was no means known by which physicists could determine the true direc- 
 tion of the current of electrons and the existence, even, of the electron 
 was not suspected. The guess which was made as to the direction of the 
 current was the opposite of the truth and it seems impossible now to 
 correct the mistake. 
 
246 SILICON, TIN AND LEAD 
 
 quadrivalent. The compound is a volatile liquid which 
 boils at 114. As iron does not combine so readily with 
 chlorine at low temperatures and ferric chloride, FeCl 3 , 
 is much less volatile, scrap tin and old tin cans are treated 
 with chlorine gas to recover the tin. The stannic chloride 
 may be easily reduced to metallic tin or may be used directly 
 as a mordant or in other ways. Stannic chloride fumes 
 in the air because it is hydrolyzed by water, giving hydro- 
 chloric acid. This shows the close relationship of tin 
 with silicon and other non-metallic elements. 
 
 Metastannic acid, H2Sn 5 Oii.9H 2 O, is an insoluble com- 
 pound formed when tin is treated with nitric acid. Its 
 formation is often used for the detection and determination 
 of tin in alloys, since all of the other metals which are 
 acted upon by nitric acid, except antimony, are converted 
 into nitrates by the acid and dissolve. 
 
 Stannic hydroxide, or stannic acid, Sn(OH) 4 , is formed 
 by precipitating a solution of stannic chloride with ammo- 
 nium hydroxide, NH^OH. It dissolves in either strong 
 acids, such as hydrochloric or sulfuric acid, or in strong 
 bases, such as sodium hydroxide. It is, therefore, both a 
 base and an acid. The solution in sodium hydroxide con- 
 tains sodium stannate, Na 2 SnO3. 
 
 Fireproofing of Cotton Goods. If a piece of cotton cloth 
 is dipped in a solution of sodium stannate, Na 2 SnO 3 , 
 and then, after squeezing out the excess of the solution and 
 drying the cloth, it is dipped in a solution of ammonium 
 sulfate and again squeezed and dried, the stannic acid, 
 H 2 SnO 3 , or stannic oxide, SnO 2 , formed by these processes 
 combines so firmly with the fiber of the cloth that no amount 
 of washing will remove it. Flannelette and other forms of 
 cotton cloth which have been treated in this manner will no 
 longer catch fire and burn. If the process could be generally 
 introduced, many fatal accidents from burning might be 
 prevented. (See Professor W. H. Perkin's address before 
 
LEAD 247 
 
 the International Congress of Applied Chemistry in 1912.) 
 In some states laws have been passed requiring the fire- 
 proofing of stage curtains in schools and public buildings. 
 
 Lead, Occurrence, Metallurgy. Lead is found in nature 
 chiefly in the form of galena, PbS, a heavy, black mineral 
 with a bright, metallic luster on fresh surfaces. The mineral 
 crystallizes and cleaves in cubes. When heated in the air it 
 is converted into a mixture of sulfate, PbSO4, and oxide, PbO. 
 When this mixture is heated with more of the original galena 
 all of the lead is reduced to the metallic state, if the pro- 
 portions are properly chosen: 
 
 PbSO 4 + PbS = 2Pb + 2SO 2 
 2PbO + PbS = 3Pb + S0 2 
 
 Properties and Uses of Lead. Lead is a soft metal easily 
 cut with a knife and yielding so readily to pressure that it 
 may be forced through a die of the proper shape into the 
 form of lead pipe. It melts easily (327) but at a consider- 
 ably higher temperature than tin or solder. It is dissolved 
 to a slight degree by pure water and is not suitable for 
 pipes which are to carry water for household use, because 
 lead compounds are very poisonous and even a very minute 
 amount of lead taken into the system daily through a series 
 of weeks or months may be dangerous. 
 
 Lead is very slightly attacked by dilute sulfuric acid, even 
 at high temperatures, and it is used for the lead chambers 
 and for the evaporating pans for the manufacture of sulfuric 
 acid. Lead dissolves, however, in hot, concentrated sulfuric 
 acid. The conduct of iron is just the reverse. It dissolves 
 easily in dilute sulfuric acid but is scarcely affected by the 
 concentrated acid. Accordingly, the concentration of the 
 dilute acid is carried on in leaden pans to the point where 
 the lead begins to dissolve and then the process is finished 
 in iron. 
 
 Lead is used for bullets and shot because it is the heaviest 
 
248 SILICON, TIN AND LEAD 
 
 of the cheap metals and because it is so easily given the 
 desired forms. Solder is an alloy of lead with tin. Babbitt 
 and antifriction metals are alloys with tin and antimony, 
 which gives hardness and sharpness of outline. Lead is 
 also a constituent of the fusible alloys used for safety plugs 
 in steam boilers, in electric circuits and elsewhere. 
 
 Oxides of Lead. Lead forms three important oxides: 
 litharge, PbO, red lead, Pb 3 O4, and lead dioxide, PbO 2 . 
 Litharge or lead monoxide is formed when lead, or galena, 
 is heated to a rather high temperature in the air. Red 
 lead, Pb 3 O4, is formed by heating litharge or white lead at a 
 somewhat lower temperature. It is to be considered as a 
 lead salt of plumbic acid, H 4 PbO4, and the formula may be 
 written Pb 2 Pb04, showing that two-thirds of the lead is in 
 the basic condition and one-third is acid. In accordance 
 with this formula, on treatment with nitric acid the basic 
 lead dissolves as lead nitrate, Pb(NO3) 2 , while the acid 
 lead remains undissolved as lead dioxide, Pb02, the anhydride 
 of plumbic acid, H4Pb04.* 
 
 Pb 2 Pb0 4 + 4HN0 3 = 2Pb(N0 3 ) 2 + Pb0 2 + 2H 2 O 
 
 Red lead is used as a pigment. It will be noticed that 
 lead dioxide is similar in formula to carbon dioxide and that 
 plumbic acid decomposes into lead dioxide and water as 
 carbonic acid, H 2 COs, decomposes into carbon dioxide and 
 water. 
 
 Storage Batteries. In a storage battery which has been 
 charged one plate contains a considerable amount of spongy 
 metallic lead and the other plate consists largely of lead 
 dioxide, Pb0 2 , in which the lead is quadrivalent. The jar of 
 the battery contains dilute sulfuric acid. It will be remem- 
 bered that some of the atoms of a plate of metallic lead in 
 contact with dilute sulfuric acid lose electrons and are 
 changed to lead ions, Pb ++ : 
 
 Pb - Pb++ + 2- 
 
STORAGE BATTERIES 249 
 
 These lead ions combine with sulfate ions, SO 4 = , of the 
 solution to give insoluble lead sulfate, PbSO 4 , but the free 
 electrons remain in the lead plate and the process is almost 
 instantly stopped by the difference in potential set up 
 between the plate and the solution: 
 
 Pb++ + SO==PbS0 4 
 
 If the lead plate is connected with the plate containing 
 lead dioxide, Pb(>2, the electrons will pass to that plate 
 and combine with the oxygen, converting it into oxygen ions, 
 = , and leaving the lead as lead ions, Pb++. 
 
 Pb0 2 + 2- = Pb++ + 20= 
 
 The lead ions will combine with sulfate ions to form lead 
 sulfate, as before, while the oxygen ions, O = , will combine 
 with hydrogen ions, H + , of the solution to form water: 
 
 0= + 2H+ = H 2 O 
 
 At the end of the discharge both plates will contain lead 
 sulfate. In charging the battery an external electromotive 
 force is applied and the lead sulfate is reduced to metallic 
 lead while the lead sulfate of the other plate is oxidized to 
 lead dioxide with the liberation of sulfate ions. The differ- 
 ence in potential between the two plates is about two volts 
 and it is evident that metallic lead and lead dioxide, contain 
 much more chemical energy than the equivalent amounts 
 of lead sulfate. 
 
 The sulfuric acid in the storage battery disappears as the 
 battery is discharged and reappears as it is charged. For 
 this reason the condition of the cell can be tested by deter- 
 mining the specific gravity of the liquid in the cell. Dilute 
 sulfuric acid is heavier than water and the density of the 
 liquid will approach that of water as the cell is discharged. 
 
 Lead nitrate, Pb(NO3)2, is an easily soluble salt which 
 can be prepared by dissolving either lead or lead oxide 
 in nitric acid. 
 
250 SILICON, TIN AND LEAD 
 
 Lead chloride, Pb.Cl 2 , requires 125 parts of cold water for 
 its solution and is formed as a precipitate when hydro- 
 chloric acid or a chloride is added to a solution of a soluble 
 lead salt. It is more easily soluble in hot water. 
 
 Lead acetate or sugar of lead, Pb(C 2 H 3 O2) 2 .3H 2 O, is 
 formed when litharge is dissolved in acetic acid. It is 
 easily soluble and has sometimes been used for hair dyes, 
 because of the black lead sulfide, PbS, formed when it is 
 applied to the hair. Such a use is, however, considered 
 dangerous. 
 
 White lead, one of the most valuable white pigments 
 that we have, is a basic carbonate of lead having approxi- 
 mately the composition 2PbCO 3 .Pb(OH) 2 . Thin discs 
 of lead are placed in earthenware pots containing acetic acid. 
 These are packed in series with layers of spent tan bark 
 or some other organic material which will ferment and 
 furnish carbon dioxide. The lead slowly corrodes and in 
 three or four months it is almost completely converted into 
 solid, brittle cakes of the basic carbonate. This is finely 
 ground and mixed with linseed oil for use as a paint. White 
 lead is very poisonous and great care is needed to protect 
 the workmen who handle it. 
 
 Chrome yellow, PbCrO 4 , is a brilliant yellow pigment 
 used, mixed with linseed oil, as a paint. It is prepared 
 by precipitating a solution of some soluble lead salt with 
 potassium dichromate, K 2 Cr 2 07. 
 
 SUMMARY 
 
 The carbon family of elements contains carbon, the most 
 important element of living bodies, and silicon, the most 
 abundant element (except oxygen) in minerals. It also 
 contains tin and lead. 
 
 The elements of the group all form dioxides, and all 
 except silicon form monoxides. 
 
SUMMARY. SILICON, TIN, LEAD 251 
 
 Silicon is found chiefly in silica and in silicates. 
 
 Carborundum is a very hard, crystalline carbide of sili- 
 con, used as an abrasive. 
 
 Silicon dioxide is found as quartz, jasper, amethyst and 
 agate. 
 
 A great variety of silicates is found in nature, especially 
 kaolin, granite, garnet, talc and asbestos. 
 
 Glass is an artificial silicate of sodium or potassium, with 
 calcium, lead or other metals. It is viscous through a wide 
 range of temperature and remains amorphous and trans- 
 parent when cold. 
 
 Soluble glass is a silicate of sodium which will dissolve 
 in water. 
 
 A colloidal solution contains particles much larger than 
 ordinary molecules but still very small. These particles 
 remain in solution either because they are too small to 
 settle out or because they cannot coalesce on account of 
 electrically charged particles which they contain. 
 
 Colloidal solutions may be separated from electrolytes 
 by dialysis. 
 
 Tin is found as cassiterite in Cornwall, England, the East 
 Indies and Tasmania. 
 
 The solution pressure of a metal is the force which causes 
 a metal to give positive ions to a solution with which it is 
 in contact and to acquire, for that reason, a negative charge. 
 
 The electromotive series is an arrangement of metals in 
 the order of the magnitude of their solution pressures. 
 
 The most important alloys of tin are solder, Babbitt metal 
 and bronze. 
 
 Tin forms two chlorides, stannous chloride and stannic 
 chloride. 
 
 Stannic hydroxide is both a base and an acid. It is used 
 to fireproof cotton goods. 
 
 Metastannic acid is formed when nitric acid acts on tin. 
 It is insoluble. 
 
252 SILICON, TIN AND LEAD 
 
 Lead is found as galena and is obtained by roasting the 
 mineral and heating the mixture of lead sulfate, lead oxide 
 and lead sulfide which is formed. 
 
 Lead is used for lead pipe, for the leaden chambers of sul- 
 furic acid works and in solder, type metal, Babbitt metal 
 and safety fuses. 
 
 Lead forms three oxides, litharge, red lead and the dioxide. 
 Red lead is lead plumbate. 
 
 Metallic lead, lead dioxide, and sulfuric acid are formed in 
 charging a storage battery; lead sulfate in discharging it. 
 
 Lead nitrate, lead chloride, and lead acetate are common 
 salts of lead. 
 
 White lead is a basic lead carbonate used as a pigment. 
 
 Lead chromate or chrome yellow is used as a pigment. 
 
 EXERCISES 
 
 1. How much silicon dioxide and how much carbon would be 
 required to prepare a ton of silicon? 
 
 2. How much of each will be required to prepare a ton of car- 
 borundum? 
 
 3. If a colloidal particle is spherical and has a diameter of 
 3 % oooooo of a millimeter and a specific gravity of 2, how many 
 of this size will be required to weigh one gram ? 
 
 4. Design a die which might be used to press lead through for 
 the manufacture of lead pipe. Draw a cross section of the top 
 and bottc n of the opening through which the lead must be pressed. 
 The formation of the tube depends on the welding of the clean 
 surfaces of lead brought together with the die. 
 
 5. If a mixture of lime, CaO, and litharge is heated in the 
 air calcium plumbate is formed. Write the equations for the 
 reactions, also the equations for the reactions which will occur 
 when calcium plumbate is treated with dilute hydrochloric acid. 
 In what proportions should the litharge and lime be mixed? 
 
 6. Concentrated hydrochloric acid converts lead dioxide to 
 lead chloride; what other products will be formed? Write the 
 equations. What other dioxide acts in a similar manner? 
 
CHAPTER XXII 
 GROUP III: BORON, ALUMINIUM 
 
 The third group of elements contains only two which are 
 used either in the metallic state or in their compounds for 
 important industrial purposes. These are boron and alu- 
 minium. Boron is distinctly non-metallic in its properties 
 and is practically used only in the form of its compounds. 
 Aluminium is used extensively as a metal and several of 
 its salts are also important. Each element is trivalent. 
 
 Boron, Occurrence. Boron is always found in nature 
 either as boric acid, H 3 B0 3 , or in the form of salts derived 
 from boric acid or from boric anhydride, B 2 3 . Boric 
 acid issues with steam from fissures in the ground in Tuscany, 
 Italy, and considerable quantities of the acid are obtained 
 from this source. Borax and other salts of the boric acids 
 are found in California, Nevada and other places in the 
 west, and nearly all mineral waters contain small amounts 
 of borax. 
 
 Borax, Na 4 B 2 O 7 .10H2O. Borax may be considered either 
 a salt of pyroboric acid, H 4 B 2 07, or as an acid salt of boric 
 acid, H 3 BOs. In the latter case we should write the formula 
 Na 2 Hio(BO 3 )4. 5H 2 O. Boric acid is so very weak an acid, 
 however, that the salt shows no acid properties in solution 
 but, on the contrary, it is hydrolyzed by water and the 
 solution has an alkaline reaction : 
 Na 2 Hi (BO 3 ) 4 + 2H+ + 20H~ = 2Na+ + 4H 3 BO 3 + 2OH- 
 
 The alkaline reaction is, of course, due to the hydroxide 
 ions,- OH~, in the solution. This property of a mild alkali 
 caused by its hydrolysis makes it suitable for washing 
 
 253 
 
254 BORON, ALUMINIUM 
 
 flannels and other delicate fabrics which are liable to injury 
 if sodium hydroxide or some other strong alkali were used. 
 
 Borax for Welding. Borax Beads. When borax is 
 heated it swells up at first and loses water but finally melts 
 to a clear glass having the composition Na 2 B 4 O7. The 
 formula may also be written Na 2 (BO 2 )2.B 2 O3. The boric 
 anhydride, B 2 3 , which it contains, may combine with 
 metallic oxides to form borates exactly as silicon dioxide, 
 SiO 2 , combines with lime, litharge or other oxides in the 
 manufacture of glass. Thus if borax is sprinkled on a piece 
 of hot iron which is covered with a coating of iron oxide 
 the boric anhydride will combine with the oxide to form 
 a fusible glass. By bringing two pieces of iron together 
 with borax between them the film of oxide on the surface 
 of each is dissolved and, on pounding, the red-hot surfaces of 
 pure iron unite to form a perfect weld, the liquid borax being- 
 forced out from between them. Hot borax glass will dis- 
 solve the oxides of many other metals and some of these 
 give characteristic colors which may be used to identify 
 compounds of the metals in the laboratory. 
 
 Boric Acid, HsBO 3 . Borax and boric acid kill or prevent 
 the growth of bacteria and this property has led to their 
 use in preserving food and as an eye-wash. The use as a 
 food preservative is now forbidden or strictly regulated by 
 federal and state laws. 
 
 Boric acid is only slightly soluble in water. It is prepared 
 by adding hydrochloric or sulfuric acid to a warm solution 
 of borax. On cooling, the boric acid crystallizes from the 
 solution. 
 
 Aluminium, Occurrence. Aluminium is the third ele- 
 ment in abundance in the crust of the earth, oxygen and 
 silicon being first and second, and iron fourth. It is chiefly 
 found in the form of silicates. A large proportion of the 
 rocks which formed the crust of the earth in the earliest 
 geological time must have consisted of silicates of alumin- 
 
ALUMINIUM 255 
 
 him, iron, calcium, sodium, potassium and other elements. 
 During many millions of years these silicates have been 
 broken down and worked over by the action of water, ice 
 and air, with the aid, in some cases, of vegetation and earth- 
 worms. During the process, partly by mechanical agencies, 
 partly by the solvent action of the water, the original 
 minerals have been ground to exceedingly fine particles 
 which have been carried away by the water and deposited 
 as clay, shales and sedimentary rocks. When the process 
 is most complete the result is kaolin, a hydrated silicate of 
 aluminium. In such cases the sodium, potassium and 
 other elements of the original rocks have been dissolved 
 out and carried away, but usually the clays and shales are 
 mixtures of kaolin with minute particles of silica, SiO 2 , 
 and of other minerals. 
 
 Aluminium is also found in the oxide, A1 2 3 , as the 
 mineral corundum. In its crude form corundum is known 
 as emery and is used in making emery wheels, emery paper, 
 etc., because it is the hardest natural mineral except the 
 diamond. For this use it has been largely replaced by car- 
 borundum, SiC, which is even harder than corundum, 
 though less hard than the diamond. Rubies are red crystals 
 of corundum colored by a little chromium. They are now 
 made artificially. Another form of the mineral having a 
 blue color, is sapphire. 
 
 Metallic aluminium is prepared by the electrolysis of 
 aluminium oxide, A1 2 O 3 , dissolved in cryolite or some other 
 mineral which melts at a low temperature and contains no 
 water. The electrolysis is carried out at such a temperature 
 that the metallic aluminium collects in the molten form in 
 the bottom of the iron pot (Fig. 45) which is used for the 
 electrolysis. It is drawn out from time to time through 
 an opening in the end, near the bottom. 
 
 The metal melts at 657, a considerably lower tempera- 
 ture than the melting points of silver and copper ^ Its 
 
256 
 
 BORON, ALUMINIUM 
 
 specific gravity is 2.6, nearly the same as that of glass and 
 only one-third the specific gravity of iron. The metal has 
 a bright, silver- white luster and does not tarnish readily. 
 For this reason and because of its lightness it is used to a 
 considerable extent for kitchen utensils. It is also used 
 for telephone wires, and in an increasing number of valuable 
 alloys with copper and other metals. Aluminium dissolves 
 
 readily in a solution 
 of sodium hydroxide, 
 giving sodium alumi- 
 nate, Na 3 AlO 3 . It 
 also dissolves in sul- 
 furic acid, giving 
 aluminium sulfate, 
 A1 2 (S0 4 ) 3 . 
 
 Goldschmidt's 
 Thermite Process. 
 When aluminium is 
 burned to the oxide, 
 
 YIG. 45. A1 2 O 3 , a large amount 
 
 of heat is liberated, 
 
 much more than by the burning of iron or other common 
 metals. If metallic aluminium is mixed with ferric oxide, 
 Fe 2 3 , and the mixture is ignited, the reaction represented 
 by the equation: 
 
 2A1 + Fe 2 3 = A1 2 O 3 + 2Fe 
 
 takes place very rapidly with the evolution of such a quan- 
 tity of heat that the whole mass is heated far above the 
 melting point of iron. The process has been used for 
 welding steel rails, repairing broken shafts and other 
 similar purposes. 
 
 Aluminium Sulfate, A1 2 (SO 4 ) 3 .18H 2 O. This salt is 
 prepared by the decomposition of clay with sulfuric acid. 
 It is easily soluble in water and is extensively used under 
 
 '.I I 
 
ALUM, EARTHENWARE 257 
 
 the name of "alum" for the clarification and purification of 
 water. It reacts with the calcium bicarbonate, CaH 2 (CO 3 ) 2 , 
 present in practically all natural waters, giving a precipi- 
 tate of insoluble aluminium hydroxide, A1(OH) 3 , which is 
 gelatinous and adheres to fine particles of clay and to 
 bacteria in the water in such a manner that they can be 
 removed along with the aluminium hydroxide bv^ filtration. 
 
 AI 2 (S0 4 ) 3 + 3CaH 2 (CO 3 ) 2 = 2A1(OH) 3 + 3CaSO 4 + 6C0 2 
 
 Alum, KA1(SO 4 ) 2 .12H 2 O, is a salt with an astringent, 
 sweetish taste, which is easily prepared by crystallizing a mix- 
 ture of aluminium sulfate and potassium sulfate. It is used 
 extensively in the cheaper grades of baking powders. 
 For this purpose it is mixed with "baking soda, " sodium 
 bicarbonate, NaHCO 3 . The alum and sodium bicarbonate 
 react in the same manner as the aluminium sulfate and 
 calcium bicarbonate in the purification of water. 
 
 Alum is also extensively used as a mordant in dyeing. 
 The aluminium hydroxide formed by its hydrolysis combines 
 with dyes to form insoluble compounds. 
 
 Brick, Earthenware, Porcelain. It has been pointed out 
 that ordinary clays are mixtures of kaolin, quartz and other 
 minerals. Some of these minerals melt at temperatures 
 which can be easily obtained in ordinary furnaces, while 
 quartz and pure kaolin require very high temperatures for 
 their fusion. When clay has been molded in the forms 
 desired for brick, tile, earthenware or the " biscuit" of 
 porcelain it may be "burnt" in such furnaces, the "burn- 
 ing" consisting in heating the mass to such a temperature 
 that a part of the minerals present melt or sinter and 
 bind the silica and other minerals, which are infusible at 
 the temperature used, into a hard solid mass. In molding 
 the articles the clay is mixed with a small quantity of water 
 which forms a colloidal solution with the very fine particles 
 of the clay and renders the mass plastic. 
 
 17 
 
258 BORON, ALUMINIUM 
 
 Materials made from clay in the manner which has been 
 described remain porous and are not suitable for domestic 
 uses. For such purposes they must be covered with a glaze 
 to make them impervious to water. The glaze is sometimes 
 formed by throwing salt into the furnace toward the close 
 of the burning. The salt volatilizes and its sodium cojp- 
 bines with the silicon and other elements of the clay to 
 form a fusible glass while the chlorine of the salt combines 
 with the hydrogen of the water present in the gases of 
 the furnace and escapes as hydrochloric acid. Porcelains 
 are usually glazed by the application of finely ground feld- 
 spar. Glazes containing lead are often used for the cheaper 
 kinds of earthenware. 
 
 SUMMARY 
 
 - Boron and aluminium are the only common elements of 
 Group III. 
 
 Boron is non-metallic, aluminium metallic. Each is 
 trivalent. 
 
 Boron is found in boric acid, borax and other borates. 
 
 Borax is a salt of pyroboric acid. It is hydrolyzed by 
 water and is used for laundry purposes, for welding of iron 
 and in the detection of elements by the colors they give to 
 borax beads. 
 
 Boric acid is a very weak acid and has valuable antiseptic 
 properties. 
 
 Aluminium is found in clay and in many silicates. It is 
 obtained by the electrolysis of a solution of aluminium oxide 
 in cryolite or some similar double fluoride. 
 
 Aluminium is a light metal somewhat resembling silver 
 in appearance. It is used for kitchen utensils, for electric 
 conductors and in many alloys. 
 
 In the thermite process the reaction between aluminium 
 and ferric oxide or some other oxide is used to obtain a very 
 high temperature for welding and repairing iron or steel. 
 
EXERCISES. ALUMINIUM 259 
 
 Alum is used in baking powders and as a mordant. 
 
 Aluminium sulfate is used as an aid in clarifying and puri- 
 fying water. 
 
 Brick, earthenware and porcelain are made by the partial 
 melting or sintering of clay or kaolin. 
 
 The glaze for earthenware and porcelain is a glass con- 
 taining feldspar or fusible silicates of sodium, lead or other 
 metals. 
 
 EXERCISES 
 
 1. In what proportions should aluminium and ferric oxide Lo 
 mixed for use in the thermite process? 
 
 2. Many alums are known in which ammonium, NH 4 , sodium, 
 Na, or some other univalent metal takes the place of potassium, 
 and others in which iron, Fe, chromium, Cr, or some other 
 trivalent metal replaces the aluminium. Write the formulas for 
 the following: 
 
 Ammonium alum containing aluminium and ammonium. 
 Ferric ammonium alum containing iron and ammonium. 
 Chrome alum containing chromium and potassium. 
 Rubidium alum containing aluminium and rubidium. 
 
 3. The best known baking powders contain either cream of 
 tartar and sodium bicarbonate mixed with starch, or alum and 
 sodium bicarbonate mixed with starch. Calculate the per- 
 centage composition of each kind of powder, assuming 40 per cent 
 of starch and 60 per cent of the other ingredients. How many 
 grams of carbon dioxide will be furnished by a pound of each 
 (1 Ib. = 453 grams)? How many liters of carbon dioxide? 
 
 4. What will be the cost per pound of each baking powder at 
 the following wholesale prices: 
 
 Starch, 2^ cents per Ib. 
 
 Alum, 2 cents per Ib. 
 
 Cream of tartar, 26 cents per Ib. 
 
 Sodium bicarbonate, 3> cents per Ib. 
 
CHAPTER XXIII 
 
 GROUP II, FIRST DIVISION: ALKALI-EARTH METALS, 
 CALCIUM, STRONTIUM, BARIUM, RADIUM 
 
 Elements of the Second Group. The second group of 
 elements falls into two very distinct divisions. Magnesium 
 and calcium, the first well-known elements of the group, are 
 rather closely related and strontium, barium and radium 
 resemble calcium. Zinc, cadmium and mercury of the sec- 
 ond division differ very markedly from the metals of the 
 first division. In spite of some resemblance to calcium, 
 magnesium is usually classified with the second division of 
 the group. 
 
 The hydroxides of the calcium division are strong bases 
 and the metals of the division are called for that reason 
 alkali-earth metals. The hydroxides of zinc and cadmium of 
 the second division are less basic and mercury, the last 
 element of the division, forms no hydroxide. 
 
 The sulfates of the alkali-earth metals are less and less 
 soluble with increasing atomic weights and the extreme 
 insolubility of radium sulfate is used in separating it from 
 other elements. 
 
 Calcium ranks fifth in abundance among the elements of 
 the earth's crust and is one of the most important of the 
 elements. Calcium phosphate, Ca 3 (PO 4 ) 2 , is the principal 
 constituent of the bony skeleton of our bodies and of 
 the bodies of all vertebrate animals. Calcium carbonate, 
 CaCO 3 , forms the skeleton of the coral insects which have 
 built the immense coral reefs in the ocean and of insects 
 which have built the limestones which were formed during 
 
 260 
 
FLUORSPAR. LIME 261 
 
 millions of years of geologic time. Some of the limestones 
 have been metamorphosed into marble by heat and pres- 
 sure. Some of them contain magnesium carbonate, MgCOs, 
 as well, and in Switzerland whole mountains are made up 
 largely of dolomite, MgC0 3 .CaCO 3 , a mineral containing 
 calcium carbonate and magnesium carbonate in nearly 
 equimolecular proportions. 
 
 Gypsum is a soft, crystalline, hydrated calcium sulfate, 
 CaSO 4 .2H 2 O, which is called alabaster in the massive 
 white forms used for vases and statuettes. A clear crystal- 
 line form which cleaves in thin sheets is known as the min- 
 eral selenite. Cruder forms are used in the manufacture of 
 plaster of Paris and as a fertilizer to furnish sulfur and cal- 
 cium to soils poor in these elements and also to retain 
 ammonia in the soil. 
 
 Fluorite or Fluorspar, CaF 2 , has been spoken of as the 
 chief source of fluorine and its compounds. It is also exten- 
 sively used as a flux in foundries because it melts at a 
 comparatively low temperature and dissolves substances 
 which otherwise might remain mixed with the iron and 
 weaken it. 
 
 Calcium is now easily obtained by the electrolysis of 
 melted, anhydrous calcium chloride, CaCl2. It is a white, 
 crystalline metal which decomposes water at ordinary 
 temperatures with the formation .of calcium hydroxide, 
 Ca(OH) 2 . It is sometimes used to remove traces of water 
 from absolute alcohol. 
 
 Calcium Oxide or Lime is manufactured on a large scale 
 by heating limestone, CaCO 3 , in a " lime-kiln. 7 ' In the older 
 'ime-kilns the pieces of limestone and coal or wood were 
 nixed and burned in some form of furnace. When the 
 materials were cool the lime was removed from the furnace 
 vid a new mixture put in. At the present time, the lime- 
 stone and fuel are charged in alternate layers into the top 
 of a cylindrical tower and the finished lime is raked out at 
 
262 ALKALI-EARTH METALS 
 
 the bottom from time to time without stopping the process. 
 This method is, of course, much more economical of fuel. 
 
 Calcium Hydroxide or Slaked Lime. Calcium oxide 
 combines with water directly with the evolution of a con- 
 siderable quantity of heat : 
 
 OH 
 CaO + H 2 O = Ca(OH) 2 or Ca/ 
 
 OH 
 
 As the lime combines with the water it falls to a very fine 
 'powder and if a considerable excess of water is used a 
 milky suspension of the slaked lime in water, known as 
 milk of lime, is obtained. If a larger amount of water is 
 used, the calcium hydroxide settles out leaving a clear 
 solution of the hydroxide, called lime water. As calcium 
 hydroxide is only slightly soluble in water the solution is 
 always very dilute. It is used in the laboratory for the 
 detection of carbon dioxide. Why? 
 
 Mortar. When milk of lime is mixed with sharp, clean 
 sand a plastic mass is obtained which is used between layers 
 of brick or stone in building walls and also for plastering the 
 walls of rooms. As the water is absorbed by the bricks or 
 dries out in the air a mass of considerable initial strength is 
 produced, but the strength is greatly increased by the slow 
 conversion of the calcium hydroxide to crystalline calcium 
 carbonate, CaC0 3 , by the carbon dioxide of the atmosphere: 
 
 Ca(OH) 2 + CO 2 = CaC0 3 + H 2 O 
 
 The water liberated keeps the air in freshly plastend 
 rooms moist for some time. The crystals of calcium rj, r - 
 bonate adhere firmly to the particles of sand and binj the 
 whole together to a solid, hard mass. 
 
 Cement is manufactured by heating to a high temperatuio 
 a finely powdered mixture of limestone, nearly free from 
 magnesium carbonate, with a clay rich in silica. The tern- 
 
HARD WATERS 263 
 
 perature required is very much higher than that of a lime- 
 kiln as the silica and alumina must be brought into combi- 
 nation as calcium silicate and aluminate and this requires 
 the sintering of the mass to a " clinker." The "clinker" 
 is again finely ground and mixed with a little plaster of 
 Paris, giving a finished cement of the following composition : 
 
 Loss on ignition 0-2 per cent 
 
 Silica, Si0 2 ! 15-20 per cent 
 
 Alumina, A1 2 3 3-8 per cent 
 
 Ferric oxide, Fe 2 3 3-6 per cent 
 
 Lime, CaO 58-64 per cent 
 
 Magnesia, MgO 0-4 per cent 
 
 Potash and soda, K 2 O, Na 2 0-2 per cent] 
 
 Sulfur trioxide, 80s 0-2 per cent 
 
 When mixed with water the cement combines with it,, 
 forming partly crystals of calcium hydroxide, Ca(OH) 2 , 
 partly hydrates of the calcium and aluminium silicates, 
 which "set" to a hard mass. Sand, gravel or other 
 materials are added to increase the volume. As the harden- 
 ing is not due to carbon dioxide, cement will set under 
 water and is often called "hydraulic cement." 
 
 Temporary and Permanent Hardness of Waters. Salts 
 of calcium or magnesium render water "hard," that is, a 
 considerable amount of soap must be used with such water 
 before it will have the soft feeling characteristic of a soapy 
 water, because the calcium and magnesium combine with 
 the fatty acids of the soap, giving a curdy precipitate, and 
 the soap cannot produce the natural effect on the water 
 till all of the calcium has been precipitated. 
 
 Many natural waters contain calcium carbonate, CaCO 3 , 
 held in solution by carbonic acid, H 2 CO 3 , as calcium bi- 
 carbonate, CaH 2 (CO 3 )2. When such a water is boiled the 
 carbonic acid of the bicarbonate dissociates into carbon 
 dioxide and water and the carbon dioxide escapes. The- 
 
264 ALKALI-EARTH METALS 
 
 calcium carbonate which remains is practically insoluble 
 in water and separates as a precipitate. When the hard- 
 ness of the water is due to calcium bicarbonate it is possible 
 to remove it by boiling the water and allowing the calcium 
 carbonate to settle, and this kind of hardness is called 
 temporary hardness. 
 
 Temporary hardness may also be removed by adding just 
 the right amount of milk of lime, Ca(OH) 2 . Why? 
 
 Calcium sulfate, CaSO 4 , is also appreciably soluble in 
 water but, as it is not held in solution by carbonic acid, 
 boiling the water for a short time will not cause the calcium 
 sulfate to precipitate. For this reason hardness due to 
 calcium sulfate is called " permanent hardness. " When 
 the water is boiled away, as is done in a steam boiler, the 
 calcium sulfate separates and produces a scale which adheres 
 to the surface within the boiler and is particularly trouble- 
 some. Permanent hardness may be removed by the ad- 
 dition of the proper amount of sodium carbonate. Why? 
 
 Calcium Sulfate, Gypsum, CaSO 4 .2H 2 O, Plaster of 
 Paris, 2CaSO 4 .H 2 O. When gypsum is heated it loses 
 part of its water and is converted into plaster of Paris, which 
 has the composition 2CaS0 4 .H 2 O. If the gypsum is 
 heated to too high a temperature all of the water may be 
 driven out but the anhydrous calcium sulfate, CaS0 4 , 
 does not combine readily with water and is not suitable 
 for the uses to which plaster of Paris is applied. When 
 plaster of Paris is mixed with water to the consistency of a 
 thick cream the mixture will set in a few minutes to a solid 
 mass of gypsum, CaSO 4 .2H 2 O, as the plaster combines 
 with the water and the mass crystallizes. The material 
 is used in making plaster casts of works of art, in the "hard 
 finish" for plastered walls, and for many other purposes. 
 
 Calcium Chloride, CaCl 2 . A solution of calcium chloride 
 is easily obtained by dissolving marble in a solution of 
 hydrochloric acid. When the solution is evaporated the 
 
CALCIUM PHOSPHATES 265 
 
 hydrates of calcium chloride which remain retain water so 
 firmly that it is necessary to heat the residue to 250 or 
 above before all of the water is expelled. The porous 
 mass which remains will absorb moisture greedily from 
 ordinary air or from moist gases and it is often used to dry 
 gases in the laboratory. If exposed to ordinary air, the 
 salt will finally absorb enough water to deliquesce, dis- 
 solving in the water which it attracts to itself. 
 
 A solution of calcium chloride, which may be cooled to 
 20-25 below C. without freezing, is now extensively 
 used in refrigerating machines to surround the cans con- 
 taining distilled water which is to be frozen to artificial ice. 
 
 OC1 
 
 Chloride of Lime, Ca<^ , has been described in con- 
 
 Cl 
 nection with hypochlorous acid (p. 86). 
 
 Calcium Phosphate, Ca 3 (PO 4 ) 2 . The occurrence of cal- 
 cium phosphate as the chief mineral constituent of bones 
 has been already mentioned. Calcium phosphate is also 
 found in large deposits in North and South Carolina, 
 Tennessee, Georgia and Florida and is extensively used 
 as one of the most important constituents of commercial 
 fertilizers which are applied to land poor in available 
 phosphorus. As calcium phosphate is almost insoluble 
 in water, the mineral is often made more easily soluble 
 and available for plants by treating it with sulfuric acid to 
 convert it into the acid phosphate, CaH 4 (PO 4 )2. Tho 
 mixture of calcium sulfate and acid calcium phosphate is 
 called, commercially, a " super-phosphate. " Write the equa- 
 tion for the reaction. Very finely ground calcium phosphate 
 is also slowly taken up by plants and during a series of 
 years may be as useful as the acid phosphate. 
 
 Calcium Carbide, CaC 2 , is manufactured by heating 
 lime and coke in electrical furnaces. It was used at first 
 for the preparation of acetylene (p. 200) for illuminating 
 
266 ALKALI-EARTH METALS 
 
 purposes, but large quantities are now converted into cal- 
 cium cyanamide for use as a fertilizer. 
 
 Calcium Cyanamide, or "Lime -nitrogen," CaCN 2 . At a 
 high temperature calcium carbide unites with nitrogen to 
 form calcium cyanamide and carbon: 
 
 CaC 2 + N 2 = CaCN 2 + C 
 
 Water converts calcium cyanamide to calcium carbonate 
 and ammonia: 
 
 CaCN 2 + 3H 2 = CaCO 3 + 2NH 3 
 
 The ammonia formed is readily available in the soil for 
 the growth of crops and for this reason "lime-nitrogen" 
 is a valuable constituent for fertilizers to be used in soils 
 that are deficient in nitrogen. 
 
 Atomic Weight of Calcium. Law of Dulong and Petit. 
 It has been pointed out (p. 133) that the most satisfactory 
 method of selecting the true atomic weight of an element 
 consists in finding the weight of the element contained 
 in a gram-molecular volume (22.4 liters at and 760 mm.) 
 of that gaseous compound which contains the smallest 
 quantity of the element in this volume. But calcium forms 
 no compound whose weight in the gaseous form has been 
 determined, and a considerable number of other elements 
 form no compounds which can be converted into gases 
 without decomposition. The atomic weights of such ele- 
 ments must, of course, be selected in a different manner. 
 For this purpose the law of Dulong and Petit, discovered 
 in 1819, has been useful. These chemists found that the 
 quantity of heat required to raise the temperature of one 
 gram atom of an element one degree is approximately 6.6 
 calories. If this quantity of heat is applied to 7 grams of 
 lithium or to 65 grams of zinc or to 200 grams of mercury 
 it will, in each case, raise the temperature one degree. 
 
 The law is also frequently stated that the specific heat 
 
LAW OF DULONG AND PETIT 
 
 267 
 
 of 'an element multiplied by its atomic weight is a constant 
 quantity. The following table will make this clear: 
 
 Element 
 
 Specific 
 heat . 
 
 Atomic 
 weight 
 
 Atomic heat 
 (sp. ht. X at. wt.) 
 
 Lithium 
 
 94 
 
 7 
 
 6 6 
 
 Graphite (at 11) 
 
 16 
 
 12 
 
 1 9 
 
 Graphite (at 977) 
 
 467 
 
 12 
 
 5 6 
 
 Silicon 
 
 16 
 
 28 4 
 
 4 5 
 
 Calcium 
 
 17- 
 
 40 
 
 6 8 
 
 Zinc 
 
 093 
 
 65 4 
 
 6 1 
 
 Bromine 
 
 084 
 
 80 
 
 6 7 
 
 Mercury 
 
 033 
 
 200 
 
 6 7 
 
 Lead 
 
 03 
 
 207 
 
 6 2 
 
 
 
 
 
 It will be seen from the table that graphite and silicon 
 depart rather widely from the law, though the former 
 approaches it more closely at high temperatures. All of 
 the metallic elements and all elements having atomic 
 weights above 40 conform approximately to the law. The 
 law is, at best, however, only approximate and is of service 
 only in selecting between rather widely divergent possible 
 values for an atomic weight. Thus the atomic weight of 
 calcium might be 20, 40 or 60, according as the formula, 
 of the chloride is CaCl, CaCl 2 or CaCl 3 . But of these three 
 values only an atomic weight of 40 agrees with the law. 
 If the atomic weight were 20, the atomic heat (see above) 
 would be 20 X 0.17 = 3.4. If it were 60, the atomic heat 
 would be 10.2. An atomic weight of 40 gives an atomic 
 heat of 40 X 0.17 = 6.8, which approximates closely to the 
 average value (6.6) for other elements. 
 
 The laws of Avogadro and of Dulong and Petit have 
 usually been considered as independent and wholly un- 
 related. l A little consideration, however, shows us that if 
 
 1 See, however, G. N.Lewis, J. Am. Chem, Soc., 29, 1165 and 151G (1907), 
 
268 ALKALI-EARTH METALS 
 
 we accept the kinetic-molecular theory this is not the 
 case. At foundation Avogadro's law depends on the fact 
 that molecules of different weights exchange energies, when 
 in collision with each other or with the walls of the con- 
 taining vessel at a given temperature, in such a manner 
 that the average value of ^mv 2 (m, mass, v, velocity) is 
 constant and is independent of the weight of the molecule. 
 The law of Dulong and Petit must depend on a similar 
 property of the atoms of the elements in the solid or liquid 
 state. 
 
 Strontium, Sr, and Barium, Ba, are found as sulfates, 
 SrSO 4 and BaSC>4, and carbonates, SrCO 3 and BaC0 3 , which 
 are all nearly insoluble in water. It will be noticed that the 
 sulfate and the carbonate are also two of the most common 
 compounds of calcium. 
 
 Barium Peroxide, BaC>2, is formed when barium oxide, 
 BaO, is heated in the air. The reaction is reversible and 
 at a higher temperature or under diminished pressure the 
 peroxide decomposes into barium oxide and oxygen. This 
 property has been much used in the preparation of oxygen 
 gas. Explain how and write the equations. 
 
 Barium Chloride, BaCl 2 .2H 2 O. Barium sulfate, BaSO 4 , 
 is almost the only salt of barium which does not dis- 
 solve in dilute hydrochloric acid. For this reason barium 
 chloride is often used in the laboratory for the detection 
 of sulfuric acid and of soluble sulfates and for the determi- 
 nation of the amount of these which may be present in a 
 mixture or solution. 
 
 Radium, Discovery. In 1896, shortly after the dis- 
 covery of Rontgen or X-rays, Henri Becquerel in Paris 
 discovered that uranium compounds will affect a photo- 
 graphic plate through a layer of black paper but that the 
 effect is cut off by metals in the same manner as the X-rays. 
 After several years of laborious, painstaking investigation 
 Monsieur and Madame Curie showed that these effects 
 
RADIOACTIVE PROPERTIES 
 
 269 
 
 are due in only a very trifling degree to pure uranium 
 and that they are caused mainly by a new element, radium, 
 which is present in very minute quantities in minerals which 
 contain compounds of uranium. 
 
 Radium and its compounds exhibit four remarkable 
 properties : 
 
 1. It affects photographic plates in a manner similar to 
 the effect of X-rays. 
 
 2. It keeps itself at a higher temperature than other ob- 
 jects around it, or, in other words, it 
 
 is constantly giving off energy in the 
 form of heat. 
 
 3. It will cause air in its neighbor- 
 hood to become a conductor of elec- 
 tricity. For instance, a gold-leaf 
 electroscope (Fig. 4^), which will 
 retain its charge for a long time in 
 ordinary air as shown by the repul- 
 sion of the leaf, will be quickly dis- 
 charged and the leaf will fall if a 
 mineral containing radium is brought 
 near to it. This has been found to 
 
 be the most accurate method of detecting and measuring 
 the amount of radium and of other radioactive elements 
 in minerals, water, air or other substances. 
 
 4o Radium, when brought near to the skin for some time,, 
 may produce severe burns, somewhat similar to sunburn, 
 and it kills bacteria very much as sunlight does. The use 
 of radium in the treatment of cancer seems closely related 
 to this property. The amount of radium to be found is so 
 small and its value for medicinal and other purposes is so 
 great that compounds of the element have been sold at a 
 price of over $100,000 a gram of radium, while gold is worth 
 only about 64 cents a gram. 
 
 5. Radium causes zinc sulfide and some other substances 
 
 FIG. 46. 
 
270 ALKALI-EARTH METALS 
 
 to phosphoresce and glow in the dark. This property 
 of radium is used in the radiolite dials of watches. 
 
 Disintegration of Elements. For a thousand years or 
 more, through the middle ages, a class of men called alchem- 
 ists sought by every 'means they could think of to find some 
 method of converting lead and other base metals into gold. 
 The 'failure of their quest for the " Philosopher's Stone' 7 
 contributed to the conclusion finally reached by chemists 
 that it is impossible to transform one chemical element 
 into another. This conclusion was universally accepted 
 at the close of the nineteenth century. In spite of this 
 belief Professor Rutherford, who was then working at Mc- 
 Gill University in Montreal, proposed the startling hy- 
 pothesis that radioactive elements disintegrate with the 
 formation of other elements. Shortly after this Soddy, 
 who began with Rutherford and who continued his studies 
 with Professor Ramsay in London, demonstrated that 
 helium is formed by the spontaneous decomposition of 
 radium. In the decomposition the helium atoms are shot 
 out with a tremendous velocity, such that they will pene- 
 trate thin layers of ordinary opaque matter and affect a 
 photographic plate beneath. They also produce the other 
 radioactive phenomena referred to in the preceding 
 paragraph. 
 
 To account for these phenomena it is supposed that the 
 atoms of the radioactive element, and probably of all other 
 elements, are composed of positively charged particles 
 around which electrons 1 are revolving with a high velocity. 
 Occasionally one of the positive particles, having the weight 
 of a helium atom, gets into an unstable position with 
 reference to the other parts of the atom and is shot out. 
 After the helium atom has escaped, the residue is no longer 
 an atom .of radium but it is now an atom of radium ema- 
 nation or niton (p. 158)! 
 
 1 Atoms of negative electricity. 
 
LIFE OF AN ELEMENT 271 
 
 It is evident from the relation of radium to helium and 
 niton that radium atoms are composite, but radium cannot 
 be considered as merely a compound of helium and niton 
 in the sense in which we speak of a compound of sodium and 
 chlorine. No means is known by which niton and helium 
 can be reunited to form radium and all three have those 
 properties by which we are accustomed to distinguish 
 elements from compounds. 
 
 Having discovered that radium atoms are composite it 
 is natural to suppose that the atoms of the other elements 
 are composite, also, and that the relations of the elements 
 in the periodic system are due to this fact. An extremely 
 interesting field for investigation has been opened up in 
 this connection and there seems to be a good probability 
 that our knowledge of the structure of atoms will develop 
 rapidly. 
 
 Life of an Element. The rate at which the atoms of dif- 
 ferent radioactive elements disintegrate varies very greatly. 
 One-half of a given quantity of radium would decompose 
 into niton and helium in about 1800 years. One-half of 
 the remainder would decompose in the next 1800 years, 
 and so on. It is evident that no definite time can be given 
 when all of the element would be disintegrated, but, by 
 common consent, the time required for the disintegration of 
 one-half of its atoms is spoken of as the " half-life period of 
 the element." Radium is formed indirectly from uranium, 
 but while the half -life period of radium is 1800 years that 
 of uranium is about 6,000,000,000 years. If radium is 
 formed only from uranium, as seems probable, it is clear 
 that the amount of radium in the world must be very small 
 in comparison with the amount of uranium. The half-life- 
 period of niton is 3.8 days and only exceedingly minute 
 quantities of that element can be obtained. A cubic 
 millimeter of niton weighs only about 0.01 milligram, but 
 with a few cubic millimeters of the gas Professor Ramsay 
 
272 
 
 ALKALI-EARTH METALS 
 
 determined its boiling point and density and showed that 
 its atomic weight is about 222. The atomic weight of 
 radium is 226.4 and that of helium is 4. 
 
 Other Radioactive Elements. In the disintegration of 
 radium, as has been stated, helium atoms are shot out and 
 the residue which remains consists of niton, an element 
 with properties widely different from radium. Niton, in 
 turn, disintegrates more than 150,000 times as fast as 
 radium. As it does so it loses another helium atom and 
 the residue, radium A, is supposed to have the atomic weight 
 of 218.4, but the element is known only by its radioactive 
 properties. Radium A also decomposes rapidly but shoots 
 out electrons instead of helium atoms. As an electron 
 weighs only one-eighteen-hundredth part as much as a 
 hydrogen atom, radium B has practically the same atomic 
 weight as radium C. 
 
 Lead is supposed to be the end of the series of products 
 formed from radium. The relation between the elements of 
 the series can be seen from the following table : 
 
 Element 
 
 Atomic 
 weight 
 
 Half-life period 
 
 Atoms expelled 
 
 Radium 
 
 226 4 
 
 1760 years 
 
 Helium 
 
 Niton 
 
 222 4 
 
 3 86 days 
 
 Helium 
 
 Radium A 
 
 218.4 
 
 3 minutes 
 
 Helium 
 
 Radium B 
 
 214.4 
 
 26 . 7 minutes 
 
 Electrons 
 
 Radium C 
 
 214.4 
 
 19 . 5 minutes 
 
 Helium and electrons 
 
 Radium D 
 Radium E 
 
 210.5 
 210.5 
 
 17.3 years 
 6 . 2 days 
 
 None 
 None 
 
 Radium E 
 
 210 5 
 
 4 8 days 
 
 Electrons 
 
 Radium F 
 Inactive product, 
 
 210.5 
 
 ofifi K: 
 
 143 days 
 
 Helium 
 
 
 
 
 
 Besides the radium series of radioactive elements, which 
 begins with uranium and probably ends with lead, there is 
 
COUNTING MOLECULES 273 
 
 a second series beginning with thorium and supposed to 
 end with bismuth, and a third series beginning with actin- 
 ium, an element known only by its radioactive properties, 
 and ending in an inactive product which cannot be identi- 
 fied. The atomic weight of no element in the series is known. 
 
 Chemists of the Bureau of Mines have recently proposed 
 the use of mesothorium of the thorium series to replace 
 radium for the dials of radiolite watches. 
 
 A very surprising result of recent investigations is the 
 fact that the density and atomic weight of lead from radio- 
 active sources are different from those of ordinary lead, al- 
 though the two kinds of lead are alike in other chemical 
 and physical properties. 
 
 Counting the Number of Molecules in One Cubic 
 Centimeter of a Gas. It has been shown that one gram of 
 
 FIG. 47. 
 
 radium gives off 158 cubic millimeters of helium gas in a 
 year. It would take a little more than six years for it to 
 give one cubic centimeter. This determination made it 
 possible for Rutherford and Geiger to count the helium 
 atoms in a very small volume of the gas and estimate the 
 number in one cubic centimeter. As the helium atom and 
 the helium molecule are identical, on the basis of Avogadro's 
 law this number is the same as the number of molecules in 
 a cubic centimeter of any other gas. 
 
 The apparatus used for the experiment is shown in Fig. 
 47. A small amount of radium was placed in A, which was 
 
 18 
 
274 ALKALI-EARTH METALS 
 
 completely exhausted. A thin screen of mica at D made 
 it possible to leave a little air in C while the helium atoms 
 shot out by the radium would pass through the mica into 
 C. E was a charged wire connected with an electrometer. 
 When a helium atom shot into C the air was ionized and 
 became a conductor for electricity. The discharge of E 
 through the air to the walls of C was immediately recorded 
 by the electrometer. The amount of radium and its distance 
 from the opening covered by the mica were so chosen that, 
 on the average, ten or eleven helium atoms passed into C 
 in one minute. The radium shot out the same number of 
 atoms in all directions, of course, and by measuring the 
 distance of the radium from the mica screen and the size 
 of the opening through which the atoms passed into C, 
 it was easy to calculate how many atoms were shot out in 
 a minute or an hour from the radium used. From this 
 number and from the time required for the amount of 
 radium used to send out one cubic centimeter of helium gas 
 the number of molecules in one cubic centimeter was calcu- 
 lated. 
 
 By placing a phosphorescent screen of zinc sulfide back 
 of the mica plate a flash was produced by each helium atom 
 which struck the sulfide. By counting the flashes the 
 number of helium atoms was determined. Both methods 
 gave results which were nearly the same as the best de- 
 terminations of the number of molecules in a cubic centi- 
 meter of a gas which have been made by other methods. 
 This number is about 27,100,000,000,000,000,000. 
 
 SUMMARY 
 
 The first division of the second group contains calcium, 
 strontium, barium and radium the alkali-earth metals. 
 
 The hydroxides are strong bases, the carbonates are 
 insoluble and the sulfates are difficultly soluble or insoluble. 
 
SUMMARY. ALKALI-EARTH METALS 275 
 
 Calcium occurs as the carbonate, sulfate, phosphate or 
 fluoride. 
 
 Lime is manufactured by heating calcium carbonate. 
 
 Calcium hydroxide is used in mortar and plaster. 
 
 Cement is a mixture or compound of lime with calcium 
 and aluminium silicates. It combines slowly with water, 
 forming calcium hydroxide and hydrates which set to a 
 hard mass. 
 
 Calcium sulfate gives "permanent" hardness, which 
 cannot be removed by boiling a natural water. Calcium 
 bicarbonate gives " temporary" hardness, which can be 
 removed by boiling. 
 
 Plaster of Paris is a hydrate of calcium sulfate which com- 
 bines readily with more water to form another hydrate that 
 sets to a hard mass. 
 
 Anhydrous calcium chloride is used as a drying agent. A 
 solution of calcium chloride with a low freezing point is 
 used in refrigerating machines. 
 
 Chloride of lime is partly calcium chloride and partly 
 calcium hypochlorite. 
 
 Calcium phosphate and "super-phosphate" are used as 
 fertilizers. 
 
 Calcium carbide is used for the preparation of acetylene, 
 also for the manufacture of calcium cyanamide, or lime- 
 nitrogen. The latter is used as a fertilizer. 
 
 Barium peroxide is used in the manufacture of oxygen 
 and of hydrogen peroxide. 
 
 Barium chloride is used for the detection and determina- 
 tion of sulfuric acid. 
 
 Radium is formed very slowly by the disintegration of 
 uranium. It disintegrates, itself, into helium and niton. 
 
 There are three series of radioactive elements: the 
 radium series beginning with uranium and ending with 
 lead ; a thorium series ending with bismuth ; and an actinium 
 series ending with an unknown, inactive element. 
 
276 ALKALI-EARTH METALS 
 
 It is proposed to use mesothorium in radiolite watches. 
 
 The " half-life period" of an element is the time required 
 for one-half of the element to disintegrate. 
 
 The number of helium atoms shot out by a small quantity 
 of radium has been counted and by this method the number 
 of molecules in a cubic centimeter of a gas has been 
 determined. 
 
 EXERCISES 
 
 1. How much limestone will be required to give a ton of lime ? 
 
 2. How much s^ked lime will a ton of lime give? 
 
 3. How much barium peroxide will be required to give one liter 
 of a 3 per cent solution of hydrogen peroxide? 
 
 4. The sulfur in 1.25 grams of coal was oxidized to a sulfate 
 and the solution was precipitated with a solution of barium 
 chloride. The barium sulfate obtained weighed 0.183 gram. 
 What per cent of sulfur did the coal contain? 
 
CHAPTER XXIV 
 
 METALLURGY AND THE PREPARATION OF COM- 
 POUNDS OF THE METALS 
 
 Metallurgy. By metallurgy is meant the preparation 
 of a metallic element in the free state, generally for practical 
 use. As the human race emerged from savagery the first 
 metals to be used were copper, silver and gold, all of which 
 are found free in nature. Tin was added to the list in 
 comparatively early times because it can easily be reduced 
 from the oxide, which is the most common ore. Bronzes 
 prepared by the reduction of mixed ores were also used. It 
 was not till a comparatively high state of civilization was 
 reached, though still in prehistoric times, that men learned 
 to reduce iron from its ores in simple furnaces somewhat 
 like a blacksmith's forge. 
 
 Iron was a comparatively rare and expensive metal 
 through the middle ages. About 1500, however, the blast 
 furnace was invented for the reduction of iron ores. Coal 
 was first used in blast furnaces in 1735 (see p. 324). The 
 rapid reduction of iron on a large scale has now made iron 
 the cheapest of all the metals and in its various forms it 
 is of greater importance than all the other metals combined. 
 
 Reduction of Ores of the Common Metals. The most 
 common and important metals are iron, copper, zinc, lead, 
 tin and antimony. All these are reduced to the metallic 
 state from the oxide by one of three methods : 
 
 1. The oxides of iron, zinc and tin are reduced by fuels 
 .containing carbon 
 
 2. Sulficles of zinc and antimony are oxidized to sulfur 
 
 277 
 
278 METALLURGY 
 
 dioxide and the oxide of the metal by heating them in the 
 air and the oxide is then reduced by a fuel as in the first 
 method. 
 
 3. Sulfides of copper and lead are partially oxidized by 
 roasting in the air and the mixture of oxide or sulfate with 
 some of the original sulfide is then heated. The two com- 
 pounds reduce each other and sulfur dioxide escapes. 
 
 Electrolytic Methods. Sir Humphrey Davy discovered 
 how to prepare potassium and sodium by means of an 
 electric current a little more than 100 years ago, but it was 
 not till near the close of the nineteenth century that methods 
 of producing electric currents on a large scale by means of 
 dynamos were developed. Since that time electricity has 
 been extensively used in the refining of copper and in the 
 manufacture of aluminium and sodium and magnesium. 
 The properties of copper are such that it may easily be re- 
 duced from the aqueous solution of its salts, but in the 
 commercial production of aluminium and sodium water is 
 excluded. 
 
 Reduction by Means of Aluminium. Still more recently, 
 the preparation of aluminium by electrolysis has made it 
 possible to use the metal for the preparation of other metals. 
 The use of a mixture of ferric oxide, Fe 2 O 3 , with aluminium 
 to obtain a high temperature has been described as Gold- 
 schmidt's thermite process. If chromic oxide, C^Os, is 
 used in place of ferric oxide, metallic chromium is produced 
 and the same method may be used for the preparation of 
 other metals. 
 
 Preparation of Compounds of the Metals. Nearly all 
 compounds of the metals are salts. With a very few ex- 
 ceptions these are prepared by reversible reactions carried 
 out in solutions. These reactions proceed chiefly in one 
 direction because of one of the three following conditions: 
 
 1. A Volatile Compound is Formed. Because the vola- 
 tile compound escapes from the mixture, the equilibrium 
 
PRECIPITATION 279 
 
 of the reversible reaction is constantly shifted toward its 
 formation : 
 
 NaCl + H 2 SO 4 * NaHSO 4 + HC1 
 
 Na 2 CO 3 + 2HC1 <= 2NaCl + H 2 C0 3 
 
 H 2 C0 3 <=* H 2 + C0 2 
 
 2. A Difficultly Soluble Compound is Formed. The 
 
 separation of such a compound as a precipitate shifts the 
 equilibrium toward its formation: 
 
 BaCl 2 + Na 2 S0 4 <= 2NaCl + BaSO 4 
 NaHS0 4 + HC1 <= H 2 SO 4 + NaCl 
 NaN0 3 + KC1 * KN0 3 + NaCl 
 
 In the first illustration the barium sulfate is so nearly 
 insoluble that the reaction is practically complete and it is 
 impossible for a solution containing sulfate ions to contain 
 more than a very minute trace of barium ions. 
 
 In the second illustration sodium chloride will separate 
 only when water is present to retain the hydrochloric acid 
 in solution and when the amounts of sodium and chlorine 
 are more than sufficient to saturate the solution with sodium 
 chloride. In both this and the third case it is evident that 
 salts which are quite soluble, as well as those which are 
 ordinarily spoken of as insoluble, may be obtained by reac- 
 tions of this type. Potassium nitrate, KNO 3 , is very easily 
 soluble in hot water while common salt, NaCl, is scarcely 
 more soluble in hot than in cold water. When a mixture of 
 sodium nitrate and potassium chloride is treated with a 
 small amount of hot water, sodium chloride, the least 
 soluble of the four salts, remains undissolved and by pouring 
 off and cooling the hot solution potassium nitrate is obtained. 
 Saltpeter for the manufacture of gunpowder is obtained in 
 this way. 
 
 3. A Compound Which Ionizes to only a Slight Extent is 
 Formed. The most common case of this class of reactions 
 
280 METALLURGY 
 
 is the formation of water by the union of hydrogen, H+, 
 and hydroxide, OH~, ions. The ionization of water: 
 
 HOH <=> H+ + OH- 
 
 takes place to such a trifling extent that it is impossible to 
 haveTmore than a very few hydrogen ions in a solution 
 containing hydroxide ions or more than a very few hydrox- 
 ide ions in a solution containing hydrogen ions: 
 
 NaOH + HN0 3 + NaN0 3 + HOH 
 
 Solubility of Salts. Salts vary greatly in solubility and 
 no satisfactory reasons can be given why some salts are 
 soluble while others are insoluble. In spite of some excep- 
 tions the following general rules are useful: 
 
 1. Practically all salts of the alkali metals, sodium, 
 potassium and ammonium, are soluble in water. 
 
 2. Nearly all salts of the strong, highly ionized, monobasic 
 and bibasic acids are soluble. This includes chlorides, 
 bromides, iodides, fluorides, chlorates and perchlorates, 
 sulfites and sulfates, nitrites and nitrates. 
 
 3. Normal salts of the weak, slightly ionized, bibasic 
 acids, carbonic, H 2 CO 3 , silicic, H 2 SiO 3 , and hydrosulfuric, 
 H 2 S, and of the tribasic acids, phosphoric, H 3 P04, arsenious, 
 H 3 AsO 3 , arsenic, HsAsO^ and boric, H 3 B0 3 , are insoluble 
 with the exception of the salts of the alkalies. 
 
 The following exceptions and modifications of these 
 rules are given for reference but should be learned rather 
 by experience in the laboratory than by memorizing them. 
 
 Potassium and ammonium chloroplatinates, K 2 PtCl 6 
 and (NH 4 ) 2 PtCl 6 , and potassium perchlorate, KC10 4 , are 
 only slightly soluble. 
 
 The following salts of monobasic and bibasic acids are 
 nearly insoluble: The chlorides, bromides and iodides of sil- 
 ver, cuprous copper, mercurous mercury and lead, AgCl, 
 AgBr, Agl, CuCl, Cul, HgCl, Hgl, PbCl 2 (slightly soluble), 
 
SUMMARY. METALLURGY 281 
 
 PbBr 2 , PbI 2 , mercuric iodide, HgI 2 , calcium fluoride, CaF 2 , 
 barium sulfite, BaSO 3 , and calcium, strontium, barium and 
 radium sulfates, CaS0 4 (slightly soluble), SrS0 4 , BaS0 4 , 
 RaSO 4 . 
 
 Sulfides of the alkali metals, sodium, potassium and 
 ammonium, and of the alkali-earth metals, calcium, barium 
 and strontium, are hydrolyzed by water to a hydrosulfide 
 and hydroxide: 
 
 Na 2 S + HOH * NaHS + NaOH 
 2CaS + 2HOH <= Ca(SH) 2 + Ca(OH) 2 
 
 The sulfides of aluminium and chromium are hydrolyzed 
 to hydroxides and hydrogen sulfide : 
 
 A1 2 S 3 + 6HOH = 2A1(OH) 3 + 3H 2 S 
 
 The sulfides of all the other common metals, except 
 magnesium, are insoluble in water. 
 
 SUMMARY 
 
 The most important ores of the common metals are oxides 
 or sulfides. 
 
 Common metals are obtained by reducing the oxide, in 
 nearly all cases. This may be preceded by roasting the 
 sulfide. 
 
 Some metals are obtained and others are purified by 
 electrolysis. 
 
 A few of the less common metals are prepared by re- 
 ducing the oxide with aluminium. 
 
 Nearly all reactions for the preparation of salts are 
 reversible and depend on the escape of a volatile compound, 
 the precipitation of an insoluble compound, or the formation 
 of a slightly ionized compound. 
 
 Salts of the alkali metals and of strong monobasic and 
 bibasic acids are usually soluble, but almost all salts of 
 weak bibasic acids and of tribasic acids are insoluble. 
 
282 METALLURGY 
 
 EXERCISES 
 
 Write the equations for the following processes: 
 
 1. The reduction of ferric oxide by carbon monoxide. 
 
 2. The roasting of antimony sulfide to the oxide and the reduc- 
 tion of the latter to metallic antimony. 
 
 3. The roasting of galena to lead oxide and lead sulfate and the 
 reduction of each by heating it with lead sulfide. 
 
 4. The reduction of chromic oxide by aluminium. 
 
CHAPTER XXV 
 
 GROUP II: SECOND DIVISION; MAGNESIUM, ZINC, 
 CADMIUM AND MERCURY 
 
 Characteristics of the Metals of the Second Division. 
 
 While magnesium resembles calcium both in its chemical 
 and physical .properties and in its occurrence in nature, zinc, 
 cadmium and mercury differ in increasing degree from stron- 
 tium, barium and radium (see p. 260). The hydroxides 
 of magnesium, zinc and cadmium are insoluble and mercury 
 forms no hydroxide, while the hydroxides of strontium and 
 barium are more soluble than that of calcium and are strong 
 bases. 
 
 Magnesium. Occurrence. The occurrence of magne- 
 sium with calcium in the mineral dolomite, MgCO 3 .CaCO 3 , 
 has been referred to. The pure carbonate, MgCO 3 , is 
 known as the mineral magnesite. Magnesium sulfate, 
 MgSO 4 .7H 2 O, is easily soluble and is found in some well- 
 known mineral waters, especially in Hunyadi water, to 
 which it gives part of its medicinal value. 
 
 Metallic Magnesium is a very light metal, having a 
 specific gravity of only 1.8. It resembles zinc in its appear- 
 ance and in some of its properties, but is much more 
 active. It decomposes boiling water with the evolution of 
 hydrogen. In the form of a ribbon it can be easily burnt 
 in air, giving a very intense white light. In the form of a 
 powder mixed with potassium chlorate, which yields oxygen- 
 readily, it is used in flash-light powders for photography. 
 This is because the intense white light given when it burns 
 contains the violet rays which affect the salts of silver used 
 in photographic plates. 
 
 283 
 
284 MAGNESIUM, ZINC, CADMIUM AND MERCURY 
 
 Compounds of Magnesium. Magnesium forms the com- 
 pounds to be expected of a bivalent metal easily soluble in 
 acids: the oxide, MgO, hydroxide, Mg(OH) 2 , chloride, 
 MgCl 2 .6H 2 O, and sulfate, MgSO 4 .7H 2 O. 
 
 Zinc. Occurrence, Metallurgy. Zinc is found chiefly 
 as the sulfide, sphalerite, ZnS, but the carbonate, ZnCOs, 
 and a hydrous silicate, ZnSiO 3 .H 2 O, are also found in suffi- 
 cient amounts to be available ores. 
 
 To obtain metallic zinc the sulfide is roasted, giving sulfur 
 dioxide, SO 2 , and zinc oxide, ZnO. The sulfur dioxide is 
 now often used for the manufacture of sulfuric acid. This 
 is done partly to prevent the escape of sulfur dioxide which 
 would injure the vegetation in the neighborhood of the 
 works. 
 
 The zinc oxide is mixed with coal and heated in an earth- 
 enware retort which has an earthenware receiver, in which 
 the zinc that distils from the retort collects. 
 
 Properties of Zinc. Alloys. Zinc is more electropositive 
 than iron (p. 243). In spite of this, it is much less affected 
 than iron by the combined action of air and water. This 
 is due in part to the fact that the ferric oxide and hydroxide, 
 which we call iron rust, does not adhere closely to the iron 
 and the difference of electrical potential between the rust 
 and the iron hastens the action of air and water. The 
 oxide or hydroxide formed on the surface of zinc, on the 
 other hand, adheres closely and forms a coating which pro- 
 tects the metal from further action. This property of 
 zinc, together with the fact that it is more electropositive 
 than iron and so protects iron which is in contact with it 
 from rusting, renders zinc more effective than tin as a pro- 
 tective coating for iron. It must be remembered, however, 
 that zinc dissolves, even in weak acids, and that zinc com- 
 pounds are poisonous. This makes iron covered with zinc, 
 commonly called "galvanized iron," unsuitable for kitchen 
 utensils. 
 
ZINC. MERCURY 285 
 
 Zinc is used as the electropositive metal, which passes 
 into solution, in all kinds of primary electrical batteries, 
 including the well-known "dry" cells. 
 
 Brass is an alloy of approximately two parts of copper 
 with one of zinc. 
 
 Zinc Sulfide, ZnS, is found in nature as the mineral sphal- 
 erite. It may be prepared as a white precipitate when 
 hydrogen sulfide is passed into a neutral or weakly acid 
 solution of a zinc salt. It is the only white metallic sulfide 
 which can be prepared in this way. 
 
 Zinc Oxide, ZnO, is prepared by burning metallic 
 zinc. It is a white powder which forms an excellent white 
 paint when mixed with linseed oil. White lead and paints 
 which contain it are blackened by hydrogen sulfide because 
 lead sulfide, PbS, is black. Since zinc sulfide is white, 
 paints containing zinc oxide are not blackened in the same 
 way and such paints are more suitable than paints contain- 
 ing white lead for use in laboratories and in situations where 
 the paint is exposed to sewage gases. 
 
 Zinc Sulfate, ZnSO 4 .7H 2 O. Vitriols. The name " vit- 
 riol " seems to have been first used for ferrous sulfate or green 
 vitriol, FeSO 4 .7H 2 O, and later for sulfuric acid, which was 
 called "oil of vitriol" because it was prepared by the dis- 
 tillation of a mixture of sulfates of iron. After that zinc 
 sulfate was called white vitriol and copper sulfate, CuSO-i.- 
 5H 2 O, blue vitriol. 
 
 Cadmium closely resembles zinc in most of its properties. 
 Its sulfide, CdS, is yellow. The metal is used in some of 
 the fusible alloys. 
 
 Mercury. Metallurgy, Properties. Mercury is found 
 to a limited extent in the free state in nature but occurs 
 chiefly as the sulfide, HgS, in the red mineral cinnabar. 
 The metal is prepared from this, either by roasting it to- 
 sulfur dioxide and metallic mercury, or by mixing it with 
 iron and distilling away the mercury. 
 
286 MAGNESIUM, ZINC, CADMIUM AND MERCURY 
 
 Mercury is a heavy, mobile liquid. It melts at nearly 40 
 below zero, the Fahrenheit and Centigrade thermometers 
 coming together at that point. It boils at about 360. 
 The specific gravity is 13.6. Because it is a liquid through 
 such a convenient range of temperatures and because its 
 rate of expansion is very uniform when it is heated, it is 
 used for thermometers. Because of its specific gravity, 
 such that a column 760 mm. or 30 inches in height 
 will balance the pressure of the atmosphere } it is used for 
 barometers. 
 
 Mercury alloys easily with gold, silver and many other 
 metals. These alloys are called amalgams. If the amount 
 of the foreign metal is small, the metal dissolves in the 
 mercury and the amalgam remains liquid, but in most 
 cases only a small per cent of the foreign metal is required 
 to give a solid amalgam. The property of amalgamating 
 with gold and silver is used in separating these metals 
 from the large quantities of other minerals with which they 
 are usually mixed. 
 
 Valence of Mercury. All of the metals of both divisions 
 of Group II except mercury form exclusively compounds 
 in which the metals are bivalent, as for instance in the 
 chlorides CaCl 2 and CdCl 2 and the sulfates BaSO 4 and 
 ZnSO 4 . Mercury, on the other hand, forms mercurous 
 compounds, such as Hg 2 O and HgCl (or Hg 2 Cl2) in which it 
 appears univalent, as well as mercuric compounds, such as 
 HgO, HgCl 2 and HgSO 4 , in which it is bivalent. 
 
 Mercuric Oxide, HgO. In the second chapter an experi- 
 ment of Lavoisier was described in which he heated some 
 mercury in contact with a limited amount of air till the 
 oxygen had been converted into the red oxide of mercury. 
 He then recovered the oxygen by heating the oxide to a 
 higher temperature, decomposing it into mercury and oxy- 
 gen. Priestly had previously obtained oxygen in England 
 by heating this same oxide. The historical interest at- 
 
CALOMEL. MERCURIC FULMINATE 287 
 
 tached to these experiments has made mercuric oxide a 
 common substance in chemical laboratories. 
 
 Mercurous Chloride, Hg 2 Cl 2 , or Calomel. Mercurous 
 chloride is prepared by subliming a mixture of mercuric 
 chloride, HgCl 2 , and mercury: 
 
 HgCl 2 + Hg = Hg 2 Cl 2 
 
 Calomel is extensively used as a medicine. A widespread 
 prejudice against its use arose from its administration 
 in large doses, sometimes causing salivation and other 
 serious injuries. It is now usually given in very small doses 
 and is mixed with sodium bicarbonate, NaHCOs, to neutral- 
 ize the acid of the gastric juice and render the calomel less 
 soluble. 
 
 Calomel is practically insoluble in dilute acids and is 
 precipitated when hydrochloric acid or a chloride is added 
 to a solution of a soluble mercurous salt. 
 
 Mercuric Chloride, HgCl 2 , or Corrosive Sublimate. 
 Mercuric chloride is moderately soluble in water and easily 
 soluble in alcohol. The best antidote is the white of an egg, 
 with which it forms an insoluble compound, but this ant\- 
 dote, to be effective, must be administered very promptly. 
 
 Mercuric chloride is a very powerful antiseptic and is 
 extensively used in antiseptic surgery. 
 
 Mercuric Fulminate, Hg(CNO) 2 , a salt of fulminic acid, 
 HCNO, is used as a primer in cartridges and in the per- 
 cussion caps used to explode dynamite. It explodes 
 easily with a blow, decomposing into carbon monoxide, 
 CO, nitrogen, N 2 , and mercury, all of which are gases at the 
 temperature of the explosion. 
 
 SUMMARY 
 
 The metals of the second division of Group II, with the 
 exception of mercury, form insoluble hydroxides. Mercury 
 forms no stable hydroxide. Those with the high atomic 
 
288 MAGNESIUM, ZINC, CADMIUM AND MERCURY 
 
 weights differ very markedly from the metals of the first 
 division. 
 
 Magnesium occurs as dolomite and magnesite. 
 
 Metallic magnesium resembles zinc but is much lighter 
 and more active. It is used in flash-light powders. 
 
 Magnesium is bivalent. The oxide, hydroxide, chloride 
 and sulfate are common compounds. 
 
 Zinc occurs as sphalerite and as the carbonate and 
 silicate. It is reduced by heating the oxide with carbon 
 and distilling the zinc. 
 
 Zinc sulfide is white. 
 
 Zinc oxide is used as a white paint, which is not darkened 
 by hydrogen sulfide. 
 
 Zinc chloride is used as a disinfectant. 
 
 Zinc sulfate is white vitriol. 
 
 Cadmium is used in fusible alloys. 
 
 Mercury occurs as cinnabar. It is obtained by roasting 
 the ore or by mixing it with iron and distilling. 
 
 It is a heavy liquid, used in barometers and thermometers. 
 It forms amalgams with many metals. The gold and silver 
 amalgams are used in the recovery of gold and silver from 
 their ores. 
 
 Mercury is bivalent in some compounds, univalent in 
 others. 
 
 Mercurous chloride or calomel is used as a medicine. 
 Mercuric chloride, or corrosive sublimate, is used as an 
 antiseptic in surgery. 
 
 Mercuric fulminate is used in primers and fulminating caps 
 to detonate nitroglycerine and other explosives. 
 
 EXERCISES 
 
 1. Write the equations for the metallurgy of zinc, starring 
 with sphalerite. 
 
 2. Barium sulfate may be reduced to the sulfide by heating 
 it with carbon. The barium sulfide will dissolve in water and the 
 solution is made to react with a solution of zinc sulfate to produce 
 
EXERCISES. MAGNESIUM, ZINC, MERCURY 289 
 
 the commercial pigment called lithopone. Write the equations 
 and explain the process. 
 
 3. What is the difference between sublimation and distillation? 
 
 4. Mercury is converted into mercuric sulfate by heating with 
 concentrated sulfuric acid, a part of the acid being reduced to 
 sulfur dioxide. Corrosive sublimate is prepared by subliming 
 a mixture of mercuric sulfate with salt. Write the equations. 
 
CHAPTER XXVI 
 
 GROUP I: FIRST DIVISION; ALKALI METALS, SODIUM, 
 POTASSIUM AND AMMONIUM. SPECTRUM ANALYSIS 
 
 Properties of the Alkali Metals. The halogens, chlorine, 
 bromine, iodine and fluorine, which are univalent toward 
 hydrogen, stand at the extreme in non -metallic properties 
 and in reactivity. In a similar manner the alkali metals, 
 lithium, sodium, potassium, rubidium and caesium, which are 
 univalent toward chlorine and other halogens, stand at the 
 extreme in metallic properties and in reactivity. Freshly 
 cut surfaces of these metals, or surfaces which are protected 
 from the action of the air and moisture by melting in an 
 evacuated tube, have a bright luster resembling that of 
 silver. The metals react violently with water, liberating 
 hydrogen and forming hydroxides which are strong bases, 
 easily soluble in water. 
 
 With a very few exceptions the salts of the alkali metals 
 are easily soluble in water. The salts are highly ionized in 
 solution and the solutions are good conductors of electricity. 
 
 Sodium. The occurrence of sodium as common salt, 
 NaCl, and the preparation and properties of metallic sodium 
 were described in an early chapter of the book and should be 
 reviewed at this point. A number of additional facts will 
 be given here. 
 
 Besides the occurrence as sodium chloride, sodium is found 
 in very many of the natural silicates and also as sodium 
 nitrate or Chile saltpeter, NaNO 3 . 
 
 Sodium Chloride, NaCl. Salt is probably the most com- 
 mon mineral constituent of our food, though only a small 
 
 290 
 
SODIUM 291 
 
 quantity is required. By the physiological processes of 
 the body it is hydrolyzed, yielding hydrochloric acid, which 
 is an essential part of the gastric juice, the digestive fluid 
 of the stomach. The sodium, on the other hand, is changed 
 to sodium bicarbonate, NaHCO 3 , or disodium phosphate, 
 Na 2 HPO 4 , which are important constituents of the blood 
 and of the alkaline fluids of the digestive tract. 
 
 Salt is the source for the manufacture of chlorine, hydro- 
 chloric acid and of practically all important compounds of 
 chlorine. It is also the ultimate source for the manufacture 
 of sodium hydroxide, sodium carbonate and all important 
 compounds of sodium except sodium nitrate. These com- 
 pounds include soap and glass. It will be seen from these 
 staternents that salt fills a unique and fundamental place 
 in our chemical industries. 
 
 Sodium Nitrate, NaNO 3 . For several generations the 
 Chile saltpeter, coming from a limited area on the west coast 
 of South America, has furnished nearly all of the nitric acid 
 and saltpeter, KNO 3 , of the world. Through these com- 
 pounds the explosives used in blasting, hunting and warfare 
 have been made. Nitric acid is also used in the manufac- 
 ture of dyes and of a number of important compounds used 
 in medicine. Sodium nitrate is also extensively used in 
 fertilizers applied to soils deficient in compounds of nitro- 
 gen. It is estimated that the present source of the mineral 
 will be exhausted within a comparatively few years. 
 Fortunately methods of making nitrates and nitric acid 
 from the nitrogen of the air have been discovered and there 
 is no probability of any serious injury either to our manu- 
 factures or to agriculture when the beds of sodium nitrate 
 in Chile are gone. 
 
 Sodium Carbonate, Sal Soda, or Washing Soda, 
 Na 2 CO 3 .10H 2 O. Le Blanc Soda Process. During the 
 nineteenth century a large part of the sodium carbonate 
 used for the manufacture of soap and glass was manu- 
 
292 ALKALI METALS 
 
 factured by a process invented by Le Blanc just before 
 the French revolution. The process consists of three 
 operations : 
 
 1. Treatment of salt with sulfuric acid and heating the 
 mixture till the hydrochloric acid is expelled: 
 
 NaCl + H 2 S0 4 = NaHSO 4 + HC1 
 Nad + NaHSQ, = Na 2 SO 4 + HC1 
 
 2. Heating the sodium sulfate with carbon and calcium 
 carbonate, to reduce the sodium sulfate to sodium sulfide 
 and convert the latter into sodium carbonate : 
 
 Na 2 SO 4 + 2C ' = Na 2 S + 2CO 2 
 Na 2 S + CaCO 3 = Na 2 CO 3 + CaS 
 
 3. Separation of the sodium carbonate from the calcium 
 sulfide by dissolving the former in water. The calcium 
 sulfide is nearly insoluble in the alkaline solution and may be 
 separated from it. The sal soda may then be obtained by 
 evaporating and cooling the clear solution. 
 
 Alkalinity of Sodium Carbonate. A solution of sodium 
 carbonate turns red litmus blue and reacts alkaline toward 
 all of the common indicators. Normal sodium salts or 
 potassium salts of many other weak acids give similar alka- 
 line solutions. We may say, superficially and with some 
 truth, that these weak acids are not able to completely neu- 
 talize strong bases, such as sodium hydroxide and potassium 
 hydroxide. The phenomena may be more fully and accu- 
 rately explained by a consideration of the equilibria rep- 
 resented in the following equations: 
 
 HOH <= H+ + OH- 
 Na 2 CO 3 <= Na+ + Na+ + C0 3 - 
 C0 3 = + H+ <= HC0 3 - 
 
 combining: 
 
 Na 2 CO 3 + HOH <= Na+ + Na+ + OH~ + HCO 3 ~ 
 
SOLVAY PROCESS 293 
 
 In pure water and in any truly neutral solution the numbers 
 of the hydrogen, H + , and hydroxide, OH~, ions are equal. 
 When sodium carbonate is dissolved in the water the car- 
 bonate ions, CO 3 = , combine with the hydrogen ions of the 
 water to form hydrocarbonate ions, HCO 3 ~, first, because 
 the hydrocarbonate ions separate to only a very trifling 
 extent into hydrogen, H+, and carbonate, CO 3 = , ions, and 
 second, because even this slight degree of separation is 
 repressed by the large number of carbonate ions which are 
 present. 
 
 Similar explanations may be given in the case of other 
 salts of weak acids. 
 
 Sodium Bicarbonate or Baking Soda, NaHCO 3 ; Ammonia- 
 Soda or Solvay Process. About 1860 Ernst Solvay of Brus- 
 sels, Belgium, began the commercial development of a process 
 which had been discovered twenty years before. In this 
 process ammonia is added to a salt brine and carbon dioxide 
 is passed into the mixture: 
 
 NH 3 + H 2 O + CO 2 = NH 4 HCO 3 
 
 Ammonium bicarbonate 
 
 NH 4 HCO 3 + NaCl <=> NH 4 C1 + NaHCO 3 
 
 Sodium bicarbonate 
 
 The process depends, of course, on the fact that the so- 
 dium bicarbonate, NaHCO 3 , is the least soluble of the four 
 compounds given in the second equation. 
 
 Ammonia and ammonium chloride are much more valu- 
 able than the equivalent amounts of sodium bicarbonate. 
 The commercial success of the process, therefore, depends 
 on recovering the ammonia and using it again indefinitely. 
 This is brought about by treating the solution from which 
 the sodium bicarbonate has been separated with slaked 
 lime: 
 
 2NH 4 C1 + Ca(OH) 2 = CaCl 2 + 2NH 3 + 2H 2 O 
 
 The ammonia is volatile and can be easily separated from 
 the calcium chloride by distillation. The ammonia-soda 
 
294 
 
 ALKALI METALS 
 
 process has been found to be more economical than the Le 
 Blanc soda process and by the close of the nineteenth 
 century the competition had almost completely driven out 
 the latter method of manufacture. 
 
 It will be noticed that the principles of solubility and vola- 
 tility (p. 278) are used in both the Le Blanc and the Solvay 
 processes. 
 
 The use of sodium bicarbonate in baking powders has 
 been referred to in a previous chapter. It is easily decom- 
 posed by heat into sodium carbonate, water and carbon 
 dioxide. 
 
 Sodium Hydroxide, NaOH. As long as sodium car- 
 bonate was prepared by the Le Blanc process or by the 
 ammonia-soda process, sodium hydroxide for the manu- 
 facture of soap and for other uses was prepared by adding 
 slaked lime, Ca(OH) 2 , to a solution of sodium carbonate. 
 Calcium carbonate, CaC0 3 , is less soluble than calcium 
 hydroxide and it is also the least soluble of the four 
 substances of the equation : 
 
 Na 2 CO 3 + Ca(OH) 2 = CaC0 3 + 2NaOH 
 
 Electrolytic Sodium Hydroxide. The development of 
 cheap sources of electrical power has lead to the invention 
 
 4- 
 
 FIG. 48. 
 
 of many different forms of apparatus for the electrolysis 
 of a solution of salt, NaCl, in such a manner as to produce 
 chlorine and sodium hydroxide, both of which are valu- 
 able and extensively used. One of the best of these is the 
 Castner-Kellner apparatus shown in Fig. 48. It consists 
 of a slate box divided into three compartments by two par- 
 
SODIUM HYDROXIDE 295 
 
 titions, which fit loosely into grooves in the bottom of the 
 box. Mercury placed on the bottom seals these, giving 
 a continuous metallic layer for the three compartments 
 but preventing a dilute solution of sodium hydroxide placed 
 in the central compartment from mixing with the brine 
 in the two side compartments. Graphite anodes are used 
 in the two side compartments and an iron cathode in the 
 central one. Chlorine is evolved from the anodes and is 
 collected and used for the manufacture of bleaching powder 
 or for some other purpose. The mercury in the two side 
 compartments is negative as compared with the graphite 
 anodes and the sodium liberated at its surface combines 
 with it to form a liquid sodium amalgam. By a slight tilt- 
 ing motion the amalgam is caused to flow alternately to 
 one side and the other and is brought into the central com- 
 partment. Here it is positive with reference to the more 
 negative cathode and the hydroxide ions brought to its 
 surface combine with the sodium of the amalgam to form 
 sodium hydroxide, while the hydrogen ions of the water 
 are discharged as free hydrogen at . the surface of the iron 
 cathode. The hydroxide solution is kept at a constant 
 concentration by introducing water at one side and remov- 
 ing some of the solution at the other. Salt is added to the 
 side compartments from time to time. 
 
 Dry sodium hydroxide is obtained from the solution by 
 evaporation but it must be heated to nearly a red heat to 
 remove the last of the water. It melts at that temperature 
 and may be cast into sticks. The pure hydroxide is a white 
 solid, which absorbs water greedily from the air and deli- 
 quesces. The solution absorbs carbon dioxide, however, 
 and the solution of sodium carbonate formed will evaporate, 
 leaving a white residue. 
 
 Sodium Oxide, Na 2 O, may be prepared by heating sodium 
 hydroxide with metallic sodium. It is very rarely prepared 
 or used. 
 
296 ALKALI METALS 
 
 Sodium Peroxide, Na 2 O 2 , is prepared by heating metallic 
 sodium in a current of air. The process is carried out in 
 shallow aluminium trays. It is now extensively manu- 
 factured and is used for the helmets used in the mine rescue 
 work and to some extent for regenerating the air of subma- 
 rine boats. In contact with carbon dioxide and moisture, 
 sodium carbonate and oxygen are formed. 
 
 With cold, dilute acids, sodium peroxide gives hydrogen 
 peroxide, H 2 O 2 , which is used to bleach silk, wool, hair and 
 other substances, which would be injured by chlorine. 
 
 Sodium Sulfite, Na 2 SOs.H 2 O, is prepared by burning 
 sulfur and passing the sulfur dioxide formed through a 
 solution of sodium carbonate. It is used as a reducing 
 agent in photographic developers. 
 
 Sodium Thiosulfate, Na 2 S 2 O 3 .sH 2 O, is prepared by dis- 
 solving sulfur in a solution of sodium sulfite, Na2SO 3 . The 
 name is given because it may be considered as sodium sulfate, 
 Na 2 S(>4, in which one atom of oxygen has been replaced 
 by an atom of sulfur. The prefix "thio" is derived from a 
 Greek word meaning sulfur. 
 
 Sodium thiosulfate is usually called " sodium hyposulfite" 
 by pharmacists and photographers. It is used in "fixing" 
 photographs by dissolving those portions of the silver 
 chloride or bromide which have not been affected by the 
 light. 
 
 Sodium Silicate, or Soluble Glass, Na 2 SiO 3 , is made by 
 melting a mixture of sand and sodium carbonate. It dis- 
 solves in water but is hydrolyzed to sodium hydroxide and 
 colloidal silicic acid, giving an alkaline solution. The solu- 
 tion is used to fireproof wood and fabrics, covering them 
 with a thin, glassy coating, which renders them less in- 
 flammable. The solution (about 10 per cent) is also used 
 in preserving eggs. A thin layer, impervious to the bac- 
 teria which would cause the decay of the eggs, is formed on 
 their surface. 
 
POTASSIUM. SOFT SOAP 297 
 
 Borax, Na 2 B 4 O7.10H2O, was described under boron 
 (p. 253). 
 
 Potassium. Occurrence. The decomposition and dis- 
 integration of natural silicates during the geologic ages, 
 with the formation of soils, clays and shales, has been spoken 
 of in connection with the occurrence of aluminium. All 
 fertile soils still contain portions of the original minerals 
 and these still yield potassium or sodium, or compounds of 
 these elements, in such a form that the potassium, especially, 
 is taken up by trees and plants growing in the soil. Soils 
 from which the potassium has been removed or depleted 
 by the growth of crops for a series of years may become in- 
 fertile for that reason. The three elements in which soils 
 are most likely to become deficient are potassium, phos- 
 phorus and nitrogen. 
 
 Wood Ashes. Soft Soap. When wood is burned, the 
 larger part of the potassium which it contains is left in the 
 form of potassium carbonate, which can be readily obtained 
 from the ashes by extracting them with water and evapo- 
 rating the solution. In wooded countries this process is 
 still sometimes carried out in a simple way by taking a bar- 
 rel, boring a hole in the bottom, covering this with hay 
 or shavings to act as a filter and filling it with wood ashes. 
 The ashes are then "leached" by pouring water on the top 
 and allowing it to run through the mass. Sometimes lime 
 is put in the bottom of the barrel to change the potassium 
 carbonate, K 2 CO 3 , to potassium hydroxide, KOH. The 
 "lye" obtained in this way is concentrated by boiling it 
 down in an iron pot, if necessary, and soft soap is made from 
 it by boiling it with the refuse grease of the household. 
 
 The process of making soap which has been described was 
 common in America a few generations ago but has now al- 
 most disappeared, partly with the destruction of our forests 
 and the passing of wood as a fuel, partly with the introduc- 
 tion of cheap methods for the manufacture of sodium car- 
 
298 ALKALI METALS 
 
 bonate and sodium hydroxide and the substitution of the 
 hard soaps made from these for the soft soaps formerly used. 
 
 Potash Salts from Germany. Compounds of potassium 
 are still required in large quantities for the manufacture of 
 saltpeter, KNO 3 , and for fertilizers to be used on soils 
 which are deficient in the element. At Stassfurt, in Ger- 
 many, there are deposits of minerals which contain potas- 
 sium chloride and for some time past this has been the 
 cheapest source of potassium compounds for the whole 
 world. During the European war this supply has been cut 
 off and a very careful search has been made for potassium 
 salts which may be obtained from other sources. Among the 
 most promising of these may be mentioned natural silicates 
 which can be decomposed by processes which have not yet 
 proved to be commercially successful. Potassium salts 
 are also found in seaweeds of the Pacific Coast and serious 
 attempts are being made to utilize these. Considerable 
 quantities of potassium compounds have been obtained 
 from the water of Searle's Lake, California, from some Ne- 
 braska lakes, and also as a by-product in the manufacture of 
 cement. It is too soon to know whether the potassium 
 compounds from any other source can compete with the 
 German supply after the war. 
 
 Metallic Potassium. In 1807 Sir Humphrey Davy 
 awakened very great interest in the scientific world by 
 applying the current from a powerful electric battery to 
 moist potassium hydroxide and preparing in this way 
 minute globules of metallic potassium. He also tried the 
 experiment with moist sodium hydroxide and obtained 
 metallic sodium. Later it was found that the metal may 
 be prepared by reducing potassium carbonate with carbon 
 and by the electrolysis of potassium chloride. 
 
 Potassium is a silver white metal which tarnishes in- 
 stantly in moist air and gives so much heat as to ignite the 
 liberated hydrogen, when it is thrown on water. The hydro- 
 
SALTPETER 299 
 
 gen burns with a violet flame characteristic of potassium 
 and its compounds when volatilized in a flame. The metal 
 is kept under kerosene to protect it from the action of the 
 air. 
 
 Potassium Oxide, K 2 O, and Potassium Hydroxide, KOH, 
 are so similar to sodium oxide and hydroxide in the methods 
 of preparation and properties that they require no separate 
 description here. 
 
 Potassium Chlorate, KC1O 3 , may be prepared by passing 
 chlorine into a warm concentrated solution of potassium 
 hydroxide. Potassium hypochlorite, KC1O, is formed at 
 first and this changes to chlorate and chloride by self- 
 oxidation. Potassium chlorate is used in the laboratory 
 for the preparation of oxygen. It is also used in medicine, 
 in flash-light powders, in matches and in some of the primers 
 used in shells for firearms. 
 
 Potassium Nitrate or Saltpeter, KNO 3 . From the 
 introduction of gunpowder into Europe, about 1300, it 
 was practically the only explosive used in hunting and war- 
 fare till near the close of the nineteenth century. The 
 saltpeter for the manufacture of gunpowder was almost 
 entirely obtained, until comparatively recent times, from 
 the decay of organic matter containing nitrogen and potas- 
 sium. Considerable quantities of saltpeter formed in this 
 way have been obtained from India. During the war of 
 1812 the United States depended largely on saltpeter from 
 the Mammoth Cave, Kentucky. During more recent times 
 the world's supply of potassium nitrate has been largely 
 manufactured by the interaction of Chile saltpeter, NaNO 3 , 
 from South America and potassium chloride from Germany. 
 With the introduction of smokeless powder and other ex- 
 plosives the demand for potassium nitrate has, of course, 
 greatly decreased and the salt is no longer an important 
 factor in warfare when other nitrates can be obtained. 
 Sodium nitrate is hygroscopic and it never successfully 
 
300 ALKALI METALS 
 
 replaced saltpeter in the old forms of gunpowder. Apart 
 from that property it would be better than potassium 
 nitrate. Why? 
 
 Saltpeter is used in the curing of meats, to which it 
 imparts a desirable red color. Taken in considerable 
 quantities it is a poison, but in small quantities it is a nor- 
 mal constituent of a number of vegetables. 
 
 Gunpowder is a mixture of about 75 parts of saltpeter, 
 13 parts of charcoal and 12 parts of sulfur. This corre- 
 sponds very nearly to the equation : 
 
 2KNO 3 -f 3C + S - K 2 S + N 2 + 3C0 2 
 
 The explosion depends on the rapid formation of a large 
 volume of gases heated to a high temperature by the burning 
 of the powder. The rate of burning is regulated by the 
 size of the grains, as these burn from the surface inward. 
 For small arms a small size of grain which burns very 
 rapidly is used. For cannon the grains must be large, 
 sometimes an inch in diameter, in order to give time for 
 the heavy ball to get started before the full force of the 
 explosive is developed. 
 
 Potassium Carbonate, K 2 CO 3 . The preparation of 
 crude potassium carbonate from wood ashes has been 
 described. It is a deliquescent salt, differing in .this pro- 
 perty from sodium carbonate. 
 
 Potassium Bicarbonate or Saleratus, KHCOs. When 
 potassium compounds were more easily obtained than those 
 of sodium this salt was used for cooking, but it has now been 
 entirely replaced by the cheaper sodium salt. 
 
 Ammonium, NH4, is not, strictly speaking, a metal, 
 still less an element, but the ammonium salts are so closely 
 similar to the salts of sodium and potassium that ammonium 
 is most conveniently classed as an alkali metal. It has 
 not been prepared in the free state but when a strong solu- 
 tion of ammonium chloride is poured over sodium amalgam 
 
AMMONIUM 301 
 
 an unstable ammonium amalgam is formed. This decom- 
 poses quickly, however, into hydrogen, ammonia, NH 3 , 
 and mercury: 
 
 NH 4 Cl + Na(Hg) = NaCl + NH 4 (Hg) 
 
 Ammonium Hydroxide, NH 4 OH. When ammonia dis- 
 solves in water two things occur: the ammonia combines 
 with the water to form ammonium hydroxide and am- 
 monium hydroxide ionizes partly to ammonium, NH 4 + y 
 and hydroxide, OH~, ions: 
 
 NH 3 + HOH <=> NH 4 OH + NH 4 + + OH~ 
 
 Both reactions are reversible and may be carried to 
 practical completion in either direction. Ammonia has a 
 much lower boiling point than water, and ammonium 
 hydroxide dissociates so easily that from concentrated 
 solutions nearly pure ammonia escapes on boiling the solu- 
 tion. This is, indeed, the most convenient method of 
 preparing the gas. Ammonia also escapes from very dilute 
 solutions and all of the ammonia will be found in the water 
 which distils first from such solutions. This property is 
 used in determining minute quantities of ammonia in the 
 analysis of drinking water. Ammonia derived from the 
 organic matter of sewage may sometimes be detected in 
 this way. 
 
 Ammonium hydroxide will neutralize strong acids, giving 
 neutral salts, and in many of its properties it closely re- 
 sembles the hydroxides of the alkali metals, though it is a 
 much weaker base. By this is meant that in correspond- 
 ing dilutions there are many more hydroxide ions in the 
 solution of sodium hydroxide than in that of ammonium 
 hydroxide. 
 
 Ammonium chloride, NH^Clj may be prepared either by 
 bringing ammonia gas and hydrochloric acid gas together 
 
302 ALKALI METALS 
 
 or by neutralizing a solution of ammonium hydroxide with 
 dilute hydrochloric acid and evaporating the solution to 
 dryness : 
 
 NH 3 + HC1 = NH 4 C1 
 NH 4 OH + HC1 = NH 4 C1 + HOH 
 
 Ammonium chloride has a sharp, salty taste. When 
 the dry salt is heated it sublimes but at the same time dis- 
 sociates into ammonia and hydrochloric acid. This has 
 been proved by the weight of the gas which is formed. 
 
 Ammonium Salts. A large number of ammonium salts 
 have been prepared. Among these may be mentioned 
 ammonium sulfide, (NH 4 ) 2 S, ammonium hydrosulfide, 
 NH 4 HS, ammonium sulfate, (NH 4 ) 2 S0 4 , ammonium nitrate, 
 NH 4 N0 3 , ammonium carbonate, (NH 4 ) 2 C0 3 , ammonium 
 bicarbonate, NH 4 HC0 3 , and ammonium chloroplatinate, 
 (NH 4 ) 2 PtCl 6 . 
 
 Colored Flames. Spectrum Analysis. If a wire which 
 has been dipped in a strong solution of salt is held in the 
 blue flame of a Bunsen burner a brilliant yellow flame is 
 produced. A crystal or solution of potassium dichromate 
 which is held near such a flame will be yellow and the 
 hand or face near such a flame will assume a peculiar, 
 sallow appearance because the flame illuminates objects 
 with a pure yellow light and none of the other colors of 
 natural objects can be brought out by illumination of this 
 kind. 
 
 If the flame is brought near the slit of a spectroscope 
 (Fig. 49) at B, on looking through the telescope D a single 
 yellow line will be seen. The exact position of the line 
 can be fixed by the scale E reflected from the face of the 
 prism A, used to refract the light and separate it into its 
 different colors. 
 
 If a clean wire is dipped in a solution of pure potassium 
 chloride and then held in the flame, a wholly different, violet 
 
SPECTRA 
 
 303 
 
 flame will be given. With the spectroscope this will give a 
 red and a violet line. Still different colors and lines will be 
 given by salts of calcium, barium and other metals. 
 
 By passing electric sparks between points of other metals, 
 which cannot be volatilized in the Bunsen flame, other 
 characteristic spectra can be obtained and all of these spectra 
 may be used to identify the metals or elements that give 
 them. 
 
 FIG. 49. 
 
 The Solar Spectrum. If the sun or diffused daylight 
 is examined with a spectroscope, a spectrum with all the 
 colors of the rainbow is seen but this is crossed by many 
 fine dark lines. These lines were observed by Fraunhofer 
 early in the nineteenth century and are still called by his 
 name, though he gave no explanation for them. About 
 1860 two Germans, Bunsen and Kirchoff, took up the study 
 of these lines and showed that the position of the lines 
 corresponds exactly with the position of the bright lines 
 
304 ALKALI METALS 
 
 in the spectra of the elements. Thus there is a dark line 
 in the solar spectrum which corresponds exactly with the 
 yellow line of the sodium spectrum. 
 
 If we consider that light is a vibrating motion sent out 
 through space by a luminous object, it follows that there 
 is some rapid, periodic motion of the sodium atoms, or 
 more likely of the electrons within the sodium atoms, which 
 sends out the waves of yellow light. If these waves can 
 be sent out by sodium atoms, we may suppose that waves 
 of exactly the same length which pass sodium atoms in 
 a flame will be absorbed by the atoms. That this is so can 
 be demonstrated by placing a brilliant light which gives 
 a continuous spectrum behind a sodium flame. The con- 
 tinuous spectrum will now be crossed by a dark line where 
 the bright sodium line was before. The sodium of the 
 flame absorbs the light of its own kind from the more 
 brilliant light behind and the small amount of sodium light 
 sent on appears dark by contrast with the more brilliant 
 light on both sides of it in the spectrum. Something of 
 the same kind doubtless occurs in the sun. The central 
 portions of the sun, which are at a very high temperature, 
 send out light of all wave lengths and would give a continuous 
 spectrum of all the primary colors, if none of the light were 
 absorbed. The atmosphere of the sun, which is also at a 
 high temperature, contains the vapors of many of the 
 elements and these absorb from the original Iigh1> the vibra- 
 tions which they themselves emit, with the result that less 
 light of those wave lengths corresponding to the vibrations 
 of the atoms goes through. This gives dark lines in the 
 solar spectrum corresponding to the bright lines of sodium, 
 iron and, in all, of about thirty elements which are found 
 on the earth. The presence of hydrogen, helium and 
 other elements has also been demonstrated in many of the 
 stars so far distant that light requires many years for its 
 passage from them to the earth. 
 
SUMMARY. ALKALI METALS 305 
 
 SUMMARY 
 
 The alkali metals stand at the extreme of metallic prop- 
 erties, as the halogens are at the extreme of the non-metallic 
 elements. They decompose water at ordinary temperatures 
 and their salts are easily soluble in aqueous solutions. 
 
 Sodium occurs chiefly as common salt but also in most 
 natural silicates and in Chile saltpeter. 
 
 Sodium chloride is very widely diffused in nature and is a 
 necessary constituent of foods. It is the source of nearly 
 all compounds of sodium used in the industries. 
 
 Chile saltpeter has been the chief source of nitric acid 
 for the manufacture of explosives and for other uses. 
 
 Sodium carbonate has been made by the Le Blanc and 
 by the ammonia-soda process. The preparation of sodium 
 hydroxide by the electrolysis of salt is taking the place of 
 these processes in considerable measure. 
 
 The Le Blanc process consists of three steps: the prepa- 
 ration of sodium sulfate from salt, reduction to sodium 
 sulfide and conversion to sodium carbonate carried out 
 simultaneously, and separation of calcium sulfide from the 
 sodium carbonate. 
 
 Sodium carbonate is hydrolyzed by water because car- 
 bonic acid is a weak acid and the bicarbonate ion, HCOa", 
 ionizes scarcely more than ^water itself. 
 
 Sodium hydroxide is prepared by treating sodium car- 
 bonate with milk of lime or by the electrolysis of a solution 
 of salt. 
 
 Sodium peroxide is prepared by heating sodium in air. 
 It is used in helmets for mine rescue work and for the prepa- 
 ration of hydrogen peroxide to use in bleaching silks and 
 wool. 
 
 Sodium sulfite is a reducing agent, used in photographic 
 developers. 
 
 Soluble glass is used as a fi reproofing agent and to preserve 
 eggs. 
 
 20 
 
306 ALKALI METALS 
 
 Borax is -used in washing and for welding.. 
 
 Potassium is found in many natural silicates, in wood 
 ashes and in mineral deposits or solutions in Germany, 
 France, California and elsewhere. It is- a necessary con- 
 stituent of fertile soils. 
 
 Soft soap was formerly made with the lye from leaching 
 wood ashes. 
 
 Metallic potassium was first prepared by Sir Humphrey 
 Davy by the electrolysis of potassium hydroxide. 
 
 Potassium oxide and hydroxide are prepared in the 
 same manner as the corresponding sodium compounds. 
 
 Potassium chlorate is used in medicine, in matches, primers 
 and flash-light powders. 
 
 Potassium nitrate is used in the manufacture of black 
 gunpowder. Smokeless powders have displaced this use 
 for warfare. 
 
 Gunpowder is a mixture of charcoal, sulfur and saltpeter. 
 
 Potassium carbonate was formerly obtained from wood 
 ashes but is now manufactured from other potassium 
 salts. 
 
 Potassium bicarbonate was formerly used in cooking, 
 under the name of saleratus, but has been replaced by " cook- 
 ing soda." 
 
 Ammonium is only known in salts and as a very dilute 
 amalgam. 
 
 Ammonium hydroxide is a weak base, partly because of 
 slight ionization and partly because of dissociation to 
 ammonia and water. 
 
 Ammonium chloride is prepared by the direct union of 
 ammonia and hydrochloric acid or by neutralizing ammo- 
 nium hydroxide with hydrochloric acid. Many other am- 
 monium salts are known. 
 
 Some elements impart characteristic colors to the flame of 
 a Bunsen burner and many others develop luminous flames 
 with the electric spark. By means of the spectroscope such 
 
.EXERCISES. ALKALI METALS 307 
 
 elements may 'be identified, even when they are present in 
 very minute quantities. 
 
 Vapors of elements absorb and diffuse light of the same 
 wave lengths as those of the light which they emit. In this 
 manner the vapors in the corona of the sun cause the dark 
 lines of the solar spectrum. 
 
 EXERCISES 
 
 1. Write the equations for the reactions between sodium oxide 
 and sodium peroxide and hydrochloric acid. 
 
 2. In the presence of a little copper oxide as a catalytic agent 
 sodium peroxide dissolves in water with the evolution of oxygen. 
 Write the equation. The fused sodium peroxide containing cop- 
 per oxide is called "oxone. " 
 
 3. Sodium thiosulfate decomposes with acids, giving sulfur 
 dioxide and free sulfur. Write the equation* 
 
 4. Write the equations for the domestic manufacture of soft 
 soap, assuming that the -grease used consists chiefly of stearin. 
 
 5. Write the equation for the preparation of potassium chlorate. 
 
 6. Write the equation for the decomposition of ammonium 
 nitrate by heat. 
 
 7. What is the weight of a gram-molecular volume of! the gases 
 obtained by heating ammonium chloride? 
 
CHAPTER XXVII 
 
 GROUP I: SECOND DIVISION; COPPER, SILVER, GOLD; 
 PHOTOGRAPHY 
 
 Contrast between the First and Second Divisions of 
 Group I. There is a very marked contrast between sodium 
 and potassium on the one hand and copper, silver and gold 
 on the other; so great, indeed, that the metals are classified 
 in the same group only on account of their valences and 
 their positions in the periodic system. 
 
 Potassium and sodium are very light metals. They are 
 extremely active, tarnishing instantly in moist air, and de- 
 composing water energetically at ordinary temperatures. 
 Their hydroxides are easily soluble and are very strong 
 bases. 
 
 Copper, silver and gold are heavy metals, having a 
 brilliant metallic luster, and only copper is affected by 
 moist air. They do not decompose water at any tempera- 
 ture and are scarcely affected even by cold dilute hydro- 
 chloric or sulfuric acid. 
 
 The closest resemblance to the first division of the group 
 is found in the chlorides, Cu 2 Cl 2 (CuCl), AgCl and AuCl 
 and in the series of oxides, Cu 2 O, Ag 2 O and Au 2 O, but the 
 oxides CuO, Ag 2 O 2 and Au 2 O 3 are also known. 
 
 Copper. Occurrence, Metallurgy. The most important 
 minerals containing copper are chalcopyrite, or copper 
 pyrites, CuFeS 2 , chalcocite, Cu 2 S, and malachite, CuCO 3 .- 
 Cu(OH) 2 . It is also found as the free metal, especially 
 in the Lake Superior region. The sulfide ores are most 
 common and from these the copper is first separated as a 
 
 308 
 
COPPER 809 
 
 mixture of cuprous sulfide, Cu2S, and ferrous sulfide, FeS, 
 which is called copper matte and which melts at a low tem- 
 perature. The copper matte is then roasted in an appa- 
 ratus similar to the Bessemer converter (p. 327), the sulfur 
 escaping as sulfur dioxide. The ferrous oxide combines 
 with fine sand or silica, which is added, to form ferrous 
 silicate and the copper is reduced to the metallic state. 
 
 The copper obtained by the process which has been out- 
 lined is quite impure, containing, usually, some gold and 
 silver and larger amounts of arsenic, lead and other metals. 
 The arsenic, especially, greatly reduces the conductivity of 
 copper for electricity and renders it unfit for many industrial 
 uses. This crude copper is refined electrolytically by sus- 
 pending plates of it in a solution of copper sulfate and pass- 
 ing an electric current from the plates to a cathode of pure 
 copper. The anode of crude copper dissolves and nearly 
 pure copper is deposited on the cathode. 
 
 Copper is a red metal, which does not tarnish in dry 
 air. It is blackened by hydrogen sulfide and when exposed 
 to the weather it becomes covered with a coating of 
 basic carbonate, which has the composition of malachite, 
 CuCO 3 .Cu(OH) 2 . Copper is scarcely attacked by cold, 
 dilute hydrochloric or sulfuric acid but dissolves easily in 
 nitric acid. 
 
 Copper is used for electrical conductors, for the sheathing 
 of ships and the manufacture of brass, bronze and other 
 alloys. 
 
 Salts of Copper. Copper forms cuprous compounds, such 
 as Cu 2 O, Cu 2 S and Cu 2 Cl2, in which it appears to be uni- 
 valent, and cupric compounds, such as CuO, CuS, CuCl 2 
 and CuS0 4 , in which it is bivalent. 
 
 Cuprous Oxide, Cu 2 O. If sodium hydroxide is added 
 to a solution of copper sulfate containing Rochelle salt 
 (sodium potassium tartrate), no precipitate of copper 
 hydroxide is formed because the copper forms a complex ion 
 
310 COPPER, SILVER, GOLD; PHOTOGRAPHY 
 
 with the tartrate and there are very few copper ions, Cu ++ , 
 in the solution. On warming the solution and adding a 
 little glucose the copper is reduced to the cuprous form and 
 separates as red cuprous oxide. This process is used as 
 a method for the detection and determination of glucose 
 (p. 223) in sugar analysis and for diagnosis. 
 
 Cuprous Chloride, Cu 2 Cl 2 , is a white compound almost 
 insoluble in water but more easily soluble in concentrated 
 hydrochloric acid. It is easily prepared by digesting copper 
 turnings with cupric chloride, CuCl 2 , and concentrated 
 hydrochloric acid. The cuprous chloride is precipitated on 
 adding water to the solution. 
 
 Copper Sulfate or Blue Vitriol, CuSO 4 .5H 2 O, is the most 
 common and best known of the salts of copper. It may be 
 prepared by dissolving copper oxide in dilute sulfuric acid 
 or by dissolving metallic copper in hot concentrated sulfuric 
 acid. It is used in the electrolytic refining of copper, in 
 electrotyping and electroplating, as a mordant in dyeing, 
 and in the gravity cells which were formerly much used for 
 telegraphic purposes. A mixture of copper sulfate, slaked 
 lime and water, called Bordeaux mixture, is used for spray- 
 ing fruit trees. 
 
 Silver. Occurrence, Metallurgy. Silver is sometimes 
 found in the free state in nature but more often it occurs as 
 the sulfide, usually associated with other sulfides, especially 
 with galena, or lead sulfide. When galena is reduced to 
 metallic lead the silver is also reduced. From the lead con- 
 taining silver the metal is recovered by melting the lead and 
 mixing it thoroughly with a comparatively small amount 
 of zinc. On allowing the melted mixture to stand, nearly 
 all of the zinc rises to the top, carrying the silver with it. 
 The zinc is then skimmed off and the alloy of silver and zinc 
 with a little lead is heated in a retort to distil away the zinc. 
 The silver and lead which remain are then heated in the air 
 to oxidize the lead, leaving almost pure silver behind. 
 
SILVER; PHOTOGRAPHY 311 
 
 Silver is a white metal which does not tarnish in dry or 
 moist air at any temperature. It is easily blackened by 
 hydrogen sulfide. It dissolves in dilute nitric acid or in hot 
 concentrated sulfuric acid, very much as copper does, but 
 it is univalent in the salts formed, while copper is bivalent. 
 
 Silver coins of the United States contain 90 per cent of 
 silver and 10 per cent of copper, the copper being added to 
 give hardness and to lessen the wear of use. 
 
 Silver Nitrate or Lunar Caustic, AgNO 3 , is prepared by 
 dissolving silver in nitric acid. The salt melts easily and 
 sticks of the fused salt are used under the name of lunar 
 caustic to cauterize wounds. The name dates from the 
 time of the alchemists when silver was associated in their 
 literature with the moon and gold with the sun. 
 
 Silver Chloride, AgCl, is formed as a curdy, white pre- 
 cipitate when a solution of salt or of hydrochloric acid 
 is added to a solution of silver nitrate. Silver bromide, 
 AgBr, is a yellowish-while salt and silver iodide, Agl, 
 is a yellow salt. Each of these may be prepared in the 
 same manner as silver chloride. Each of these salts is 
 sensitive to the action of light, being reduced to metallic 
 silver or to a salt containing less of the halogen. 
 
 Silver chloride dissolves easily in ammonia and each of the 
 halogen salts dissolves in a solution of sodium thiosulfate 
 ("hyposulfite"), Na 2 S 2 O 3 . 
 
 Photography. The sensitive character of the silver halides 
 when exposed to the light is used in photography. "Dry 
 plates" are thin sheets of celluloid or plates of glass which 
 have been covered with a thin coating of an emulsion of 
 silver bromide, AgBr, in a solution of gelatin. The plate is 
 exposed in a camera for a moment to the image of the object 
 which is to be photographed. After this exposure the plate 
 appears to the eye entirely unchanged, but in some manner, 
 at present only very vaguely understood, some of the mole- 
 cules of the silver bromide are changed by the light and 
 
312 COPPER, SILVER, GOLD; PHOTOGRAPHY 
 
 are sensitive to the action of a reducing agent called a 
 "developer." After "developing" the picture, or simul- 
 taneously with the development, the portions of the silver 
 bromide which have not been affected are dissolved and 
 removed by a solution of sodium thiosulfate ("hypo- 
 sulfite" or "hypo"). This last process is called fixing. 
 Without this treatment the whole plate would become 
 black. 
 
 In the picture obtained in this way the lights and shadows 
 are, of course, reversed. The portions of the object which 
 were light appear dark in the "negative." A "positive" 
 picture is obtained by placing the negative over a sheet of 
 sensitive paper and exposing it to the light. The picture 
 printed in this manner is fixed, as before, by soaking the 
 paper in a solution of sodium thiosulfate to remove the 
 unchanged silver salt. 
 
 Gold. Occurrence, Metallurgy. Gold is found almost 
 exclusively in the free state. Occasionally large nuggets 
 weighing many pounds and worth thousands of dollars have 
 been found, but such nuggets are extremely rare. Usually 
 gold is found in small grains mixed with sand or gravel or 
 disseminated in quartz, pyrite and other rocks and minerals. 
 There is a very minute quantity of gold, worth, perhaps, two 
 or three cents, in a ton of sea water, but no profitable method 
 of recovering it has been discovered. 
 
 Gold has a specific gravity of 19.26 and because it is so 
 heavy the comparatively light minerals composing sand 
 or gravel can be separated from it easily in a current of 
 water. For rich sands and gravels the process may be 
 carried out by hand, with a pan, from which comes the 
 expression "to pan out." It is also carried out on a large 
 scale in "hydraulic mining." 
 
 For massive rocks the cyanide process is now most com- 
 mon. The ore is first broken to a powder in the stamping 
 mills and the gold and silver are extracted by a solution of 
 
GOLD 313 
 
 potassium or sodium cyanide. Oxygen from the air or some 
 oxidizing agent is required in the solution of the gold : 
 
 4Au + 8KCN + O 2 + 2H 2 O = 4KAu(CN) 2 + 4KOH 
 
 From the double cyanide of potassium and gold the latter 
 is easily precipitated by metallic zinc. 
 
 Native gold and the gold obtained by the cyanide or 
 other processes practically always contains silver. To 
 separate the two metals enough silver is added, if necessary, 
 to give an alloy containing not more than one-third of 
 its weight in gold. After melting the mixture the alloy 
 is treated with nitric acid or with hot, concentrated sulfuric 
 acid, either of which dissolves the silver, leaving the gold 
 nearly pure. 
 
 English gold coins are 22 carats fine, i.e., 2 %4 pure gold. 
 The United States coin is 900 fine, containing 900 parts of 
 gold to 100 parts of copper. 
 
 Gold does not dissolve in either of the three common 
 acids, hydrochloric, nitric or sulfuric, but it dissolves readily 
 in aqua regia. 
 
 Gold forms three oxides, Au 2 O, AuO and Au 2 3 ; and 
 three chlorides, AuCl, AuCl 2 and AuCl 3 . 
 
 SUMMARY 
 
 The second division of Group I contains copper, silver 
 and gold. These do not decompose water at any tempera- 
 ture and differ very greatly from sodium and potassium. 
 They are univalent in only a part of their compounds. 
 
 Copper is found as chalcopyrite, chalcocite and malachite. 
 The sulfide ores are concentrated to copper matte and the 
 latter is reduced by heating it with sand in a blast of air. 
 
 Copper dissolves in nitric acid or in hot, concentrated 
 sulfuric acid, but it is not affected by hydrochloric acid or 
 dilute sulfuric acid in the absence of air. 
 
314 COPPER, SILVER, GOLD; PHOTOGRAPHY 
 
 Copper forms two oxides, two chlorides and two sulfides. 
 The sulfate, blue vitriol, is the most common salt. 
 
 Red cuprous oxide is formed by the reduction of a solution 
 of a copper salt in an alkaline tart rate by means of glucose. 
 
 Cuprous chloride is prepared by reducing cupric chloride 
 by means of copper in the presence of concentrated hydro- 
 chloric acid. It is almost insoluble in water or dilute acids. 
 
 Blue vitriol is prepared by dissolving copper oxide in 
 dilute sulfuric acid. 
 
 Silver is separated from lead by adding zinc to the melted 
 metal, nearly all of the silver dissolving in the zinc and rising 
 to the top. 
 
 Silver resembles copper in its conduct toward nitric or 
 sulfuric acid. 
 
 Silver coins are an alloy of silver and copper. 
 
 Silver nitrate or lunar caustic is prepared by dissolving 
 silver in nitric acid. 
 
 Silver chloride, silver bromide and silver iodide are 
 insoluble. 
 
 In photographic dry plates silver bromide is so affected by 
 light that it is easily reduced to metallic silver. The un- 
 changed silver bromide is removed by a solution of sodium 
 thiosulfate. 
 
 Gold occurs in the free state and is obtained by washing 
 away lighter minerals with water or by dissolving it in a 
 solution of potassium cyanide and precipitating it with zinc. 
 
 Gold coins are alloys of gold and copper. 
 
 Gold does not dissolve in any of the common acids. It 
 dissolves in aqua regia. 
 
 Gold forms three oxides and three chlorides. 
 
 EXERCISES 
 
 1. Write the equation for the reduction of copper matte (p. 309) . 
 Ferrous silicate is Fe 2 Si0 4 . 
 
 2. Write the equations for the solution of copper and of silver 
 in nitric acid. Nitric oxide is evolved. 
 
EXERCISES. COPPER, SILVER, GOLD 315 
 
 3. Write the equations for the solution of copper and of silver 
 in concentrated sulfuric acid. Sulfur dioxide is formed. 
 
 4. How much silver and how much concentrated sulfuric acid 
 would be required to give 22.4 liters of sulfur dioxide? One liter 
 of sulfur dioxide weighs 2.93 grams. 
 
 6. How much chloroauric acid, HAuCl-j, can be prepared from 
 10 grams of gold? 
 
 
CHAPTER XXVIII 
 
 GROUP VI. SECOND DIVISION: CHROMIUM, TUNG- 
 STEN, URANIUM 
 
 Classification of Chromium, Tungsten and Uranium. 
 The three metals considered in this chapter belong to the 
 second division of Group VI. The compounds potassium 
 chromate, K 2 C O 4 , and sodium tungstate, Na 2 WO 4 , show 
 a close relationship to potassium sulfate, K 2 SO 4 , but while 
 there is this relationship in some of their compounds the 
 free elements show very strong contrasts. Sulfur is a solid 
 which melts at a temperature only a little above the boiling 
 point of water. It is a typical non-metallic element, both 
 in the free state and in its compounds. Chromium, tung- 
 sten and uranium, on the other hand, are metals with very 
 high melting points. 
 
 Chromium. Occurrence, Metallurgy. Chromium is 
 found chiefly in nature as the mineral chromite or chrome 
 iron ore, FeCr 2 O 4 . The mineral is isomorphous with 
 magnetite, Fe->O 4 . In other words, the two minerals have 
 the same crystalline form, indicating a similar chemical 
 structure. 
 
 Metallic chromium is prepared by igniting a mixture of 
 chromic oxide and aluminium : 
 
 Cr 2 3 + 2A1 = 2Cr + A1 2 3 
 
 Pure chromium is a very hard, brittle, crystalline metal. 
 It is used as an addition to steel to give it extreme hardness 
 and toughness for use in the armor plate of battleships. 
 
 316 
 
CHROMIUM 317 
 
 Potassium Chromate, K 2 CrO 4 , and Potassium Bichro- 
 mate, K 2 Cr 2 O 7 . When chrome iron ore is heated in the air 
 with potassium carbonate or potassium hydroxide it is 
 slowly decomposed and oxidized, the iron giving ferric oxide, 
 Fe 2 3 , and the chromium forming potassium chromate, 
 K 2 CrO 4 . Potassium chromate is very easily soluble in 
 water. The salt and its solution have a clear, lemon- 
 yellow color. Acids convert it into potassium dichromate, 
 K 2 Cr 2 O 7 , a reddish-orange salt which is much less soluble 
 than potassium chromate and which crystallizes easily. 
 
 Potassium dichromate is the best known compound of 
 chromium. It is used as the starting point for the prepa- 
 ration of nearly all other compounds of the element. It is 
 also used in acid solutions as an oxidizing agent. It is an 
 important mordant for dyeing, and is employed in chrome 
 tanning. 
 
 Chromic Anhydride, CrO 3 . When concentrated sulfuric 
 acid is added to a strong solution of potassium dichromate, 
 chromic anhydride, which is not very soluble in the con- 
 centrated acid, separates in red needles. The compound 
 corresponds to sulfuric anhydride, SO 3 . It is a vigorous 
 oxidizing agent. 
 
 Lead Chromate or Chrome Yellow, PbCrO 4 , is a yellow, 
 insoluble salt formed by adding either a solution of potas- 
 sium chromate, K 2 Cr04, or one of potassium dichromate, 
 K 2 Cr 2 7 , to a solution of sugar of lead, Pb(C 2 H 3 O 2 ) 2 . It 
 forms an excellent yellow paint. 
 
 Chrome Alum, KCr(SO 4 ) 2 .12H 2 O. If alcohol, C 2 H 6 O, 
 is added to a solution of potassium dichromate and sulfuric 
 acid it is oxidized to aldehyde, C 2 H 4 O, while the chromium 
 and potassium combine with the sulfuric acid. Chrome 
 alum, which is rather easily soluble, may be obtained by 
 evaporating the solution. The salt is isomorphous with 
 ordinary alum, indicating a similar arrangement of the 
 atoms in the two kinds of crystals. 
 
318 CHROMIUM, TUNGSTEN, URANIUM 
 
 Chromic Hydroxide, Cr(OH) 3 , and Chromic Oxide, 
 Cr 2 O 3 . The addition of ammonium hydroxide, NH 4 OH, 
 to a solution of chrome alum or of almost any chromic 
 salt produces a light green precipitate of chromic hydrox- 
 ide, Cr(OH) 3 . If the precipitate is separated and heated, 
 it decomposes into water and chromic oxide, C^Os, which 
 forms a dark green powder. 
 
 Chrome Tanning. The older methods of tanning leather 
 require the treatment of hides with the solution of tannin 
 extracted from oak or hemlock bark or obtained from other 
 sources. The process is slow and tedious, lasting for several 
 months. Another method of tanning has been developed 
 which consists in treating the hides first with potassium 
 dichromate and dilute sulfuric acid, and then with acid 
 sodium sulfite, NaHSO 3 , which causes the precipitation of 
 chromic hydroxide in the fiber. 
 
 Tungsten, The use of metallic tungsten for the filaments 
 of electric light bulbs has made the name of the element 
 almost a household word. The discovery of this use was 
 made by a chemist who was Jed to it from noticing the posi- 
 tion of the element in the periodic table. 
 
 Tungsten is also used in the high-speed tool steels. 
 Ordinary steel loses its temper and becomes soft when it is 
 heated and such steel cannot be used successfully in rapid 
 lathe work when the cutting tool often becomes very hot. 
 By adding tungsten, tools are made which retain their tem- 
 per at comparatively high temperatures. It is not too much 
 to say that this discovery has been worth many millions of 
 dollars in the saving of time for expensive machinery and 
 skilled workmen. 
 
 SUMMARY 
 
 Chromium, tungsten and uranium, of the second divi- 
 sion of Group VI, are metals with high melting points and 
 both in the free state and in those compounds in which 
 
SUMMARY. CHROMIUM 319 
 
 they have a valence of two they are distinctly metallic, 
 differing very markedly from oxygen and sulfur, of the first 
 division of the same group. 
 
 Chromium occurs as chromite. It is reduced from its 
 oxide by means of aluminium. 
 
 Chrome steel is used for armor plate. 
 
 Potassium chromate is prepared by heating chromite 
 with potassium hydroxide or carbonate in contact with air. 
 Acids change it to potassium dichromate. 
 
 Potassium dichromate is used with sulfuric acid as an 
 oxidizing agent. It is also used as a mordant in dyeing. 
 
 Lead chromate, or chrome yellow, is used as a paint. 
 
 Chrome alum is formed by the reduction of potassium 
 dichromate by alcohol in the presence of sulfuric acid. 
 
 Chrome green is a mixture of chrome yellow and Prussian 
 blue. 
 
 Hides are tanned by treating them first with potassium 
 dichromate and sulfuric acid and then with acid sodium 
 sulfite. 
 
 Tungsten is used in the filaments for electric light bulbs 
 and in the steel tools for high-speed lathes. 
 
 EXERCISES 
 
 Write equations for the following: 
 
 1. Potassium chromate from chrome iron ore. 
 
 2. Preparation of chrome alum. 
 
 3. Preparation of chrome yellow. 
 
 4. Chrome tanning. 
 
CHAPTER XXIX 
 GROUP VII. SECOND DIVISION: MANGANESE 
 
 Group VII. Manganese is the only element belonging 
 to the second division of Group VII which has thus far 
 been discovered, unless some of the radioactive elements 
 with very brief life periods belong here. We should expect 
 elements with atomic weights of about 100, 146 and 190. 
 There should also be a halogen belonging to the first division 
 of the group and having an atomic weight of 220. It 
 seems possible that atoms of Group VII with these atomic 
 weights are unstable because of the structure of such atoms. 
 According to the electron theory, halogen atoms and man- 
 ganese have seven movable electrons which they lose by 
 transfer to the oxygen atoms in such compounds as HC1O4, 
 KMnO 4 and Mn 2 O 7 , and the loss of many electrons may 
 render the atoms of higher atomic weight than manganese 
 and iodine unstable. Such an explanation is, however, 
 scarcely more than an interesting speculation at the present 
 time. 
 
 Manganese resembles chlorine in the compounds MnO 2 , 
 HMn0 4 and Mn 2 O 7 . In the compounds in which it has a 
 valence of two or three it is more metallic in character and 
 is much more closely related to iron than to the halogens. 
 
 Manganese. Occurrence, Uses. Manganese is found in 
 its purer ores chiefly as manganese dioxide, Mn0 2 . Practi- 
 cally all iron ores contain some manganese, though it is 
 usually present in only small amounts. During the first half 
 of the nineteenth century manganese came into extensive 
 use in the manufacture of chlorine for bleaching purposes. 
 As the ore became scarce and expensive, manufacturers 
 
 320 
 
COMPOUNDS OF MAGNESIA 321 
 
 of chlorine introduced methods for the recovery of the 
 manganese and conversion into a form which could be 
 used over and over in the process. In this manner the 
 demand for manganese dioxide was greatly reduced. During 
 comparatively recent times the electrolytic processes for 
 making chlorine are rapidly displacing this use of com- 
 pounds of manganese altogether. 
 
 Shortly after the middle of the nineteenth century Besse- 
 mer invented a process for manufacturing steel which makes 
 use of a form of cast iron called spiegeleisen (German for 
 " mirror iron"), because of its brilliant white luster. This 
 form of iron usually contains eight or ten per cent of man- 
 ganese. Another alloy of iron and manganese, called ferro- 
 manganese and containing, sometimes, seventy-five per 
 cent or more of the metal, is extensively used in castings for 
 car wheels and in the manufacture of other products made of 
 iron and steel. These newer uses have created a great 
 demand for ores of manganese, and iron manufacturers are 
 searching the world over for new sources of supply. 
 
 Spiegeleisen and ferromanganese are made in blast fur- 
 naces by the same processes that are used in making iron 
 (p. 324). 
 
 Oxides of Manganese. There are no less than five 
 different oxides of manganese. The best known and most 
 important is manganese dioxide, MnO 2 , sometimes called 
 black oxide of manganese because of its intense black color. 
 Its use in the manufacture of chlorine has been referred to 
 above. It is also used as the depolarizing agent in dry 
 batteries. 
 
 Potassium Manganate, K 2 MnO 4 , and Potassium Per- 
 manganate, KMnO 4 . If a mixture of manganese dioxide 
 and potassium carbonate is heated in the air a green mass 
 containing potassium manganate is formed, much as potas- 
 sium chromate is formed from chrome iron ore. The 
 oxidation may be hastened by using potassium nitrate or 
 
322 MANGANESE 
 
 potassium chlorate as an oxidizing agent. The green mass 
 dissolves in water to a green solution but if this is boiled, 
 especially if it is made slightly acid, a part of the manga- 
 nese is reduced to manganese dioxide while another part is 
 oxidized to potassium permanganate, which gives a solu- 
 tion with an intense red color. It will be recalled that 
 several compounds of chlorine show a similar self oxidation 
 and reduction. 
 
 Potassium permanganate is used as an oxidizing agent in 
 the laboratory. It is also a very efficient germicide and 
 disinfectant. If quickly applied to the wound, it is the 
 best antidote known for snake bite. 
 
 SUMMARY 
 
 Manganese is the only known element of the second 
 division of Group VII. It resembles chlorine in some of 
 its compounds but it also resembles iron. 
 
 Manganese occurs as the dioxide and as a constituent of 
 nearly all iron ores. Its alloys with iron are extensively 
 used in the manufacture of iron and steel. 
 
 Manganese dioxide was formerly used in the manufacture 
 of chlorine. It is used as a depolarizer in dry cells. 
 
 Potassium permanganate is used as a disinfectant and as 
 an oxidizing agent. 
 
 EXERCISES 
 
 Write the equations for the following: 
 
 1. Potassium manganate from manganese dioxide, potassium 
 carbonate and air. 
 
 2. Potassium permanganate, manganese dioxide and potassium 
 hydroxide from potassium manganate and water. 
 
 3. Manganese heptoxide from potassium permanganate and 
 sulfuric acid. 
 
 4. How much potassium permanganate must be dissolved in a 
 liter of water to give a solution of which 1 cc. will yield 8 mg. 
 of oxygen when the potassium permanganate is reduced to man- 
 ganese sulfate, MnS0 4 , in the presence of sulfuric acid? 
 
CHAPTER XXX 
 GROUP VIII: IRON, COBALT, NICKEL, PLATINUM 
 
 Relation of Group VIII to Other Elements. In the midst 
 of each of the longer periods of the periodic system there are, 
 in each case, three elements between the two divisions of the 
 shorter periods. Thus between the short periods , potassium- 
 manganese and copper-bromine, we have the very important 
 elements iron, cobalt and nickel. Of the six other metals 
 of Group VIII, platinum is the most interesting and im- 
 portant. All of the metals of the group have high melting 
 points above 1500. The metals in each set of three are 
 closely related to each other in physical and chemical prop- 
 erties. Thus iron, cobalt and nickel are magnetic and easily 
 soluble in the ordinary acids, especially nitric acid. The 
 platinum metals do not dissolve in nitric or hydrochloric 
 acid but are soluble in aqua regia. Valences of two, three 
 and four are the most common in the group but higher 
 valences are also found. 
 
 Iron. Importance, History. Iron was not discovered till 
 comparatively late in the history of our race, though still 
 before the beginning of written history. During long ages 
 other metals were more used than iron, but for several 
 centuries past iron has been more important than all other 
 metals put together. This importance is due chiefly to 
 three causes: 
 
 1. Iron ores are found in large quantities in many different 
 parts of the world and are easily mined and cheap. 
 
 2. Iron ores may be reduced to metallic iron easily and 
 cheaply and on a gigantic scale in the modern blast furnaces. 
 
 323 
 
324 IRON, COBALT NICKEL, PLATINUM 
 
 3. Small amounts of carbon, silicon and other elements 
 in iron, and different methods of treating the metal, render 
 it hard or soft, brittle or malleable, and give products which 
 are suitable for a great variety of uses. 
 
 Occurrence and Metallurgy of Iron. The ores of iron 
 which are used in its manufacture are all oxides, or com- 
 pounds which are converted into oxides by heat. The most 
 important are hematite, Fe 2 O 3 , limonite, Fe 2 O3.Fe 2 (OH) 6 , 
 magnetite, Fe 3 4 , and siderite, FeCO 3 . The last is a con- 
 stituent of "clay iron stone" which has been a very im- 
 portant ore in England, the country which manufactured 
 more iron than all the rest of the world during a large 
 part of the nineteenth century. 
 
 Metallic iron is found in many meteors and there is some 
 probability that there is a large amount of iron in the center 
 of the earth. 
 
 Before the days of written history men discovered how 
 to reduce iron ores and make a sort of wrought iron or 
 steel in furnaces constructed somewhat after the principle 
 of a blacksmith's forge. Simple methods of this sort have 
 continued in use among primitive people in India and Africa 
 up to modern times. The amount of iron manufactured 
 in this way was always small and during the Middle Ages 
 iron and steel were comparatively scarce and expensive. 
 
 Blast Furnaces. About 1500 a new method of manu- 
 facture similar to the modern blast furnace was invented, 
 but no one knows the name of the inventor. The old 
 processes gave a malleable iron or steel which was forged 
 while hot but which was not melted. 
 
 The blast furnace gives a hard, brittle, easily fusible 
 iron containing carbon and called pig iron. This is used 
 for castings and as the starting point for the manufacture 
 of all forms of iron and steel. 
 
 In using the blast furnace (Fig. 50) a mixture of coke, 
 iron ore and limestone is charged into the furnace at short 
 
BLAST FURNACE 
 
 325 
 
 Intervals, through the top, while a strong blast of air is 
 blown in at the bottom. As the materials make their way 
 downward the coke is 
 finally completely burned 
 at the bottom of the 
 furnace, giving at the high 
 temperature and in the re- 
 ducing atmosphere of this 
 part of the furnace carbon 
 monoxide, CO, and nitro- 
 gen. Higher up in the 
 furnace the carbon monox- 
 ide reduces the iron oxide 
 to metallic iron : 
 
 Fe 2 3 -f3CO <=* 2Fe+3CO 2 
 
 At the bottom of the 
 furnace, where the com- 
 bustion of the coke is com- 
 pleted and the temperature 
 is highest, the iron com- 
 bines with carbon and 
 silicon and melts to a 
 liquid, which is drawn off 
 from the hearth from time 
 to time. The impurities 
 of the ore, which usually 
 consist of silica, SiO 2 , and 
 a mixture of silicates, com- 
 bine with the calcium oxide 
 from the limestone to form 
 a liquid slag that floats on 
 top of the melted iron and can be drawn off separately. The 
 process goes on continuously, often for years at a time. The 
 product of the furnace is known as pig iron or cast iron It 
 
 FIG. 50. 
 
326 IRON, COBALT, NICKEL, PLATINUM 
 
 may be used directly for making stoves and for many pur- 
 poses where the iron may be given the desired form by 
 casting it in molds made of sand. Cast iron is brittle and 
 cannot be forged or welded. 
 
 Wrought Iron. In 1784 Henry Cort, in England, invented 
 a method of heating iron and stirring it in a current of 
 air in a reverberatory furnace. This was called the pud- 
 dling process. The cast iron is melted and stirred till the 
 carbon, silicon, sulfur and phosphorus have been burned 
 out, leaving a pasty mass of nearly pure iron, which can be 
 rolled into sheets, bars or rods and afterward forged or 
 welded into almost any form that is desired. This process 
 was used for 70 or 80 years in manufacturing the iron used 
 in making nails and for a great variety of other purposes. 
 Since the invention of the process for making Bessemer 
 steel, in 1856, the puddling process has been gradually 
 displaced by other methods which are more rapid and less 
 laborious. 
 
 Steel. The most important property of steel is its ability 
 to take a " temper." If heated to a moderately high 
 temperature and suddenly cooled by quenching it in water, 
 it becomes extremely hard. Under these conditions the 
 carbon is uniformly combined with or dissolved in the 
 whole mass of the iron. If the steel is heated and cooled 
 slowly, the carbon seems to combine with a small part 
 (about Y^) of the iron to form a definite compound, iron 
 carbide, Fe 3 C, while the remainder of the iron is free from 
 carbon and is soft and malleable. Intermediate forms of 
 steel with varying degrees of hardness and brittleness may 
 be obtained by heating hardened steel at carefully regulated 
 temperatures. 
 
 Bessemer Steel. In 1856 Henry Bessemer, in England, 
 described a process for manufacturing steel by blowing a 
 very strong current of air through an apparatus (Fig. 51) 
 containing several tons of cast iron. The oxygen of the air 
 
STEEL 
 
 327 
 
 burns the carbon and silicon, of the iron, and their combus-. 
 tion gives so much heat tkat the iron remains liquid after 
 these elements have been removed. Wtyen the combustion 
 of the carbon and silicon is complete the flame at the mouth 
 of the Bessemer converter drops and some* spiegeleisen is; 
 added to furnish the amount of qarbpn_ required in the 4 
 
 FIG. 51, 
 
 finished steel. After mixing the contents of the converter 
 by blowing air through it again for a very short time the 
 steel is poured into ingot molds. When it has solidified 
 the ingots are taken to the rolls while still hot and made at 
 once into rails for the railways or into other forms of mer- 
 chantable steel. 
 
 Open-hearth Steel. The original Bessemer process is 
 not adapted to pig iron containing more than a very small 
 
328 IRON, COBALT, NICKEL, PLATINUM 
 
 ii 
 
COMPOUNDS OF IRON 329 
 
 amount of phosphorus and for that reason and others a 
 very different process called the open-hearth process has 
 grown rapidly in favor during recent years. In this process- 
 a mixture of pig iron, ore, steel scraps and lime or other 
 fluxing materials is heated with gas or oil in a furnace con- 
 nected with regenerative chambers filled with a checker- 
 work of bricks (Fig. 50). These chambers are designed 
 to absorb heat from the gases coming from the furnace 
 and afterward to give up the heat absorbed, to air and gas 
 which are entering the furnace. In this manner a very high 
 temperature is secured with a comparatively small amount 
 of fuel. 
 
 By the Bessemer and open-hearth processes iron and steel 
 of almost any desired degree of hardness and strength can 
 be manufactured and the older processes of making wrought 
 iron and steel have almost completely disappeared. The 
 nails and sheet iron made by these processes rust much more 
 rapidly than those made from the old puddled iron. It has 
 been discovered, however, that the addition of a small 
 amount of copper (0.2 per cent or less) greatly lessens this 
 tendency to corrosion. 
 
 Compounds of Iron. Iron forms two series of salts, 
 ferrous salts and ferric salts. In the ferrous salts such as 
 ferrous sulfate, FeS0 4 , and ferrous chloride, FeCl 2 , the 
 iron appears bivalent, while in ferric salts, as ferric sulfate, 
 Fe 2 (SO 4 ) 3 , and ferric chloride, FeCla, the iron appears 
 trivalent. 
 
 Ferrous Sulfate or Copperas, FeSO 4 .7H 2 O, is easily pre- 
 pared by dissolving iron or ferrous sulfide in dilute sulfuric 
 acid. It is sometimes called green vitriol. 
 
 Ferrous Hydroxide is a white precipitate formed by 
 adding sodium hydroxide to a solution of ferrous sulfate 
 entirely free from oxygen or a ferric salt. It is rapidly 
 oxidized by exposure to the air, turning green and then dark 
 and finally becoming reddish-brown ferric hydroxide. 
 
330 IRON, COBALT, NICKEL, PLATINUM 
 
 Ferric Sulfate, Fe 2 (S0 4 )3. Ferrous sulfate may be 
 easily oxidized to ferric sulfate by a great variety of oxi- 
 dizing agents in the presence of sulfuric acid. It gives a 
 reddish brown precipitate of ferric hydroxide, Fe(OH) 3 , 
 when sodium hydroxide is added to its solution. 
 
 Ferric Chloride, FeCl 3 . Anhydrous ferric chloride is 
 prepared by heating iron filings or turnings in a current 
 of chlorine. It forms green scales which dissolve readily 
 in water to a yellow or reddish-yellow solution. The solu- 
 tion has an acid reaction because the salt is hydrolyzed 
 by water: 
 
 FeCl 3 + 3HOH <=> Fe(OH) 3 + 3HC1 
 
 If a solution is evaporated, a large part of the hydrochloric 
 acid escapes and anhydrous ferric chloride cannot be pre- 
 pared by this method. 
 
 Magnetic Oxide of Iron, Fe 3 O 4 . When iron is burned 
 in oxygen or when steam is passed over heated iron the 
 magnetic oxide is formed. It is formed as a closely adherent 
 black coating on the surface of red-hot iron exposed to the 
 air and the black color of the oxide is more familiar to many 
 people than the true color of iron. The magnetic oxide may 
 rust to ferric oxide or hydroxide under the action of air and 
 water but it does so less easily than a surface of metallic 
 iron. The mineral magnetite has the same composition 
 as the magnetic oxide and is one of the valuable ores of 
 iron. The magnetic oxide and magnetite are attracted by a 
 magnet, as the names imply. Other common compounds 
 of iron are only slightly magnetic. The mineral is some- 
 times found in pieces which are permanently magnetic and 
 it is then called lodestone. 
 
 Nickel is a white metal which resembles iron in some of its 
 properties. It takes a bright polish and does not tarnish 
 easily. For this reason it is deposited electrolytically as 
 "nickel plate" on iron and steel, for use on stoves, bicycle 
 
SYMPATHETIC INK 331 
 
 handles and for a great variety of practical and ornamental 
 purposes. 
 
 Nickel silver is an alloy of nickel, copper and zinc used as a 
 basis for silver-plated ware. 
 
 Nichrome is an alloy of chromium and nickel which has a 
 high melting point and which is proving very useful in 
 laboratories for wire used in electrical resistance furnaces, 
 for thermocouples, for triangles to support crucibles and 
 for many other purposes. 
 
 Nickel five-cent pieces are made of an alloy containing 
 75 per cent of copper and 25 per cent of nickel. 
 
 Nickel is bivalent in its more common compounds, such 
 as nickel sulfate, NiSO 4 , and nickel chloride, NiCl 2 . 
 
 Cobalt is another metal resembling iron. It is very little 
 used in the metallic state but cobalt oxide gives an intense 
 blue color to glass and is much used for that purpose. 
 Cobalt gives pink solutions which are complementary to 
 the green solutions of nickel salts. Mixtures of the two 
 solutions may be nearly colorless. 
 
 Anhydrous cobalt chloride is green, however. As anhy- 
 drous cobalt chloride absorbs water from moist air but 
 loses it easily on warming, cobalt chloride is used as a 
 sympathetic ink. 
 
 If the trunk and limbs of a tree are drawn on paper with a 
 lead pencil and the leaves are sketched with a solution of 
 cobalt chloride, the leaves will be nearly or quite invisible 
 in moist air but will appear on warming the paper. 
 
 Platinum. A large part of the platinum of the world 
 has come from the Ural Mountains in Russia. Some of it 
 comes from Brazil and Colombia, in South America, and 
 small quantities from California, British Columbia and 
 Alaska. The metal is very valuable for laboratory uses 
 because of its high melting point and because it does not 
 tarnish in air at any temperature and does not dissolve 
 in any single acid or base in common use. Its coefficient 
 
332 IRON, COBALT, NICKEL, PLATINUM 
 
 of expansion is so nearly the same as that of glass that 
 platinum wires may be sealed in the walls of eudiometers 
 and other forms of laboratory apparatus, forming a gas- 
 tight joint which does not crack or leak. For this reason 
 platinum has been much used for the leading-in wires of 
 electric light bulbs. 
 
 Platinum has been used in Russia for coins, but the 
 fluctuation in value in modern times has made such use 
 undesirable. 
 
 The very great increase in the price of platinum in recent 
 years has been due to its use in jewelry. As platinum is 
 very important as a catalyst in the manufacture of sulfur 
 trioxide, and for the oxidation of ammonia to nitric acid, 
 the use of the metal for articles of jewelry should be 
 discontinued. 
 
 Chloroplatinic Acid. Potassium Chloroplatinate. Plati- 
 num dissolves in aqua regia, giving a solution of chloro- 
 platinic acid, H 2 PtCl 6 , which has often been called "platinic 
 chloride." The latter name is given correctly only to the 
 compound PtCl 4 , which may be prepared by heating chloro- 
 platinic acid in a current of chlorine. 
 
 If potassium chloride KC1, is added to a solution of 
 chloroplatinic acid, a precipitate of potassium chloroplati- 
 nate, K 2 PtCl 6 , separates. This compound is used for the 
 detection and quantitative determination of potassium. 
 Ammonium salts give a similar precipitate of ammonium 
 chloroplatinate, (NH 4 ) 2 PtCl 6 . 
 
 SUMMARY 
 
 The nine elements of Group VIII are situated in groups 
 of three in the midst of the long periods, between the first, 
 and second divisions of each period. The elements vary 
 in valence, with valences of two, three and four most 
 common. 
 
SUMMARY. IRON, ETC. 333 
 
 Iron is the most important metal. The manufacture of 
 articles of iron and steel was carried on before the beginning 
 of written history. Its manufacture on a very large scale 
 is comparatively modern. 
 
 Pig iron or cast iron is made in blast furnaces. It con- 
 tains carbon and silicon as necessary ingredients, and sulfur 
 and phosphorus as impurities. 
 
 Wrought iron and steel were made in ancient times 
 directly from the ores. From near the close of the 
 eighteenth till after the middle of the nineteenth century 
 wrought iron was made by the puddling process. 
 
 The process for Bessemer steel was invented about the 
 ' middle of the nineteenth century, that for open-hearth steeT 
 some years later. In all of the processes the carbon, silicon, 
 phosphorus and sulfur are more or less completely burned 
 out of the iron and iron or steel containing a very small 
 amount of carbon is produced. 
 
 The hardness of steel is dependent on the amount and the 
 form of the carbon which it contains. 
 
 Copperas is prepared by dissolving iron in dilute sulfuric 
 acid. 
 
 Ferrous hydroxide is prepared by precipitation. It is 
 white but turns dark in the air. 
 
 Ferric sulfate is prepared by oxidizing copperas in the 
 presence of sulfuric acid. 
 
 Anhydrous ferric chloride is prepared by heating iron in 
 chlorine. It is hydrolyzed by water. 
 
 Magnetic oxide of iron is formed by heating iron to a high 
 temperature in the air. 
 
 Nickel resembles iron but oxidizes less easily. 
 
 Nickel silver contains copper, nickel and zinc. 
 
 Nichrome is an alloy of nickel and chromium. It has a 
 very high melting point. 
 
 Nickel is bivalent. 
 
 Cobalt also resembles iron. 
 
334 IRON, COBALT, NICKEL, PLATINUM 
 
 Sympathetic ink faiay be made from a solution of cobalt 
 chloride. 
 
 Platinum is found native. It is used for crucibles and as 
 a catalyst in the manufacture of sulfuric acid and of nitric 
 acid. 
 
 The most common laboratory compounds of platinum are 
 'chloroplatihic acid and potassium chloroplatinate. 
 
 EXERCISES 
 
 1. One hundred parts of water dissolve 48 parts of copperas at 
 20. What strength of dilute sulfuric acid must be used to dis- 
 ; solve ferrous sulfide and give a solution of copperas just saturated 
 at that temperature? 
 
 2. If sulfuric acid of 20 per cent is saturated with ferrous sulfate 
 by dissolving iron in it, what per cent of the copperas formed 
 Should separate on cooling the solution to 20? 
 
 3. If platinum is worth $1.75 per gram, how much is a gram of 
 chloroplatinic acid worth, not taking account of the labor and 
 materials other than the platinum used in its preparation? 
 
 4. How much pig iron containing 2.5 per cent of carbon, 2.0 
 per cent of silicon and 0.5 per cent of other elements besides iron 
 can be made from one ton of an iron ore containing 60 per cent 
 of iron? 
 
 5. Write the equation for the oxidation of ferrous sulfate to 
 ferric sulfate by potassium permanganate in the presence of 
 sulfuric acid. 
 
 6. How many milligrams of iron in the form of ferrous sulfate 
 will be oxidized by 1 cc. of the solution referred to in 4, page 322? 
 
CHAPTER XXXI 
 ANALYSIS 
 
 Qualitative analysis has for its object the determination 
 of whether certain elements, or certain groups of elements, 
 such as the ammonium group, NH 4 , the sulfate group, 
 SO 4 , the nitrate group, NOs, and others, are present in a 
 given substance or mixture. In most cases an element is 
 separated by converting it into some well-known and easily 
 recognizable compound. Thus hydrogen may be recog- 
 nized by converting it into water, or silver by converting it 
 into silver chloride. 
 
 It would be impossible within the scope of this text-book 
 to give full directions for the qualitative analysis of mix- 
 tures containing all of the more common elements and 
 radicals. The following experiments will, however, illus- 
 trate the methods which are used in analysis. 
 
 Separation of Lead, Silver and Mercurous Mercury, 
 Metals Whose Chlorides are Nearly Insoluble. Try the 
 following experiments with 3 cc. of a solution of lead 
 nitrate. Add some dilute hydrochloric acid, allow the pre- 
 cipitate to settle, pour off the solution, add 5 cc. of water, 
 allow to settle, pour off and repeat a second time. Add to 
 the precipitate 5 cc. of water, boil and pour off the solution 
 into a clean test-tube and notice that the lead chloride which 
 dissolves crystallizes from the solution on cooling. Repeat 
 the treatment with hot water till all is dissolved. Add a 
 drop of dilute sulfuric acid to the solution of lead chloride. 
 
 Repeat the same experiments with a solution of silver ni- 
 trate. When convinced that silver chloride is* practically 
 
 335 
 
336 ANALYSIS 
 
 insoluble in hot water, add some ammonia to the precipi- 
 tate. Test the solution by adding nitric acid. 
 
 Repeat the experiments with a solution of mercurous 
 nitrate. 
 
 Record the results obtained and by studying them devise 
 a method for separating and detecting the three metals, 
 when all are present. Prepare a solution containing the 
 metals and demonstrate their separation. 
 
 Separation of Mercuric Mercury, Lead, Copper and 
 Bismuth, Metals Whose Sulfides are Insoluble in Dilute 
 Acids. Take 3 cc. of a solution of mercuric chloride and 
 pass hydrogen sulfide through the solution. Filter on a 
 small filter supported in a funnel. Wash the precipitate 
 five or six times by filling the filter with water and allowing 
 the water to run through. Rinse the precipitate back into 
 the test-tube, allow it to settle and pour off the water. 
 Boil the precipitate with dilute nitric acid, then add some 
 hydrochloric acid and boil again. To the solution contain- 
 ing mercuric chloride add a solution of stannous chloride, 
 SnCl 2 . The precipitate is either mercurous chloride or 
 metallic mercury, according to the amount of stannous 
 chloride used. 
 
 Repeat the first part of the above experiments, using a 
 solution of lead nitrate. The lead sulfide will dissolve in 
 the nitric acid. Add to the solution some dilute sulfuric 
 acid, pour the solution into a porcelain dish and evaporate 
 nearly to dryness but not after fumes of sulfuric acid appear. 
 When cold rinse back into a test-tube and notice the pre- 
 cipitate of lead sulfate. 
 
 Repeat the experiments again with solutions of copper 
 sulfate and of bismuth chloride. After the evaporation 
 with sulfuric acid add ammonia to the solution. Filter 
 off the precipitate of bismuth hydroxide, BiOOH, dissolve 
 it on the filter by dropping dilute hydrochloric acid over 
 it and add a considerable amount of water to the solution 
 
QUANTITATIVE ANALYSIS 337 
 
 of bismuth trichloride, Bids. The precipitate is bismuth 
 oxychloride, BiOCl. 
 
 Devise a method of separating the four metals and apply 
 it to a solution containing all four. 
 
 Detection of Sulfates, Chlorides and Nitrates. Test 
 solutions of a number of the common salts of the laboratory 
 with barium chloride followed by hydrochloric acid, with 
 silver nitrate followed by nitric acid and with copper 
 turnings and sulfuric acid. More sensitive tests for 
 nitric acifa are described in books on qualitative analysis 
 and also further details about the tests for sulfates and 
 chlorides. 
 
 Quantitative analysis is designed to furnish information as 
 to the exact quantity of substances which are present in a 
 compound or mixture. For this purpose the element or 
 group is usually converted into some insoluble compound 
 which may be collected on a filter and washed free from 
 all other substances. The filter is then burned in a porcelain 
 or platinum crucible and the compound is weighed. Special 
 precautions are often required to avoid reduction of the 
 compound by the carbon of the filter or for other reasons. 
 
 The determination of silver or of chlorine as silver chloride 
 and the determination of barium or of the sulfate ion, SO 4 , 
 as barium sulfate may be given as illustrations of such a 
 process. 
 

 
INDEX 
 
 Absolute temperature, explana- 
 tion, 34, 40. 
 
 Absolute zero, explanation, 34. 
 
 Acetaldehyde, as absorbent of 
 acetylene, 201. 
 
 Acetic acid, in vinegar, 20; 
 preparation, 226, 235. 
 
 Acetone, as absorbent of acety- 
 lene, 201. 
 
 Acetylene, preparation ; proper- 
 ties; use as an illuminating 
 gas, 200, 208; blowpipe, use 
 of, 26 ; burners ; generators, 201 ; 
 decomposition by heat, 201; 
 why flame gives intense light; 
 why explosive; when non- 
 explosive, 201. 
 
 Acid calcium phosphate, prepara- 
 tion, 171, 265. 
 
 Acid calcium sulfite, preparation, 
 use, 117, 118. 
 
 Acid potassium tartrate, pre- 
 paration, properties; structure, 
 use, 228; action in baking 
 powders, 229 
 
 Acid salts, how formed, 77, 79. 
 
 Acid sodium carbonate, action 
 in baking powders, 229 
 
 Acid sodium sulfate, prepara- 
 tion, 80; why so named, 80; 
 by-product in preparation of 
 nitric acid, 144. 
 
 Acid sodium sulfite, preparation; 
 properties, use as germicide; 
 for preparation of sulfur diox- 
 ide, 117, 118; use in tanning, 
 318. 
 
 Acid sulfates, properties, 125. 
 
 Acid taste, indicates hydrogen, 
 20. 
 
 Acids, litmus test; sour taste, 
 11; properties, 73; mono- 
 basic, dibasic, tribasic, defi- 
 nition, 76, 118; bibasic, form 
 acid salts, 76; modern defi- 
 nition, 78; strength of in 
 relation to solubility of sul- 
 fides, 115. 
 
 Actinium, series of derivatives, 
 273. 
 
 Agate, form of silicon dioxide, 
 238. 
 
 Air, relation to combustion, 6; 
 determination of weight of 
 liter of, 39; weight of gram- 
 molecular volume, 137; com- 
 position of, 155; proof of mix- 
 ture in, weights of the four 
 gases, 159, 161. 
 
 Alabaster, properties, use, 110, 
 261. 
 
 Albumin, occurrence, 230, 235. 
 
 Alcohol, preparation; distillation, 
 225; source of ethylene, 199; 
 use; denatured, 226. 
 
 Aldehyde, by-product in prepa- 
 ration of chrome alum, 317. 
 
 Alfalfa, fixation of nitrogen by, 
 138. 
 
 Alizarin, from coal tar, 233. 
 
 Alkali-earth metals, why so 
 named ; properties ; sulfates, 
 properties, 260; list of; prop- 
 erties of compounds of, 274. 
 
 339 
 
340 
 
 INDEX 
 
 Alkali metals, properties of salts 
 of, 217; solubility of salts 
 of, 280; list of; properties; 
 salts of, properties, 290, 305. 
 
 Alkaloids, occurrence; properties; 
 list of, 232, 236. 
 
 Allotropic forms, of elements, 
 16, 17; oxygen, 15, 111; sulfur, 
 111; of phosphorus, 171; of 
 carbon, 187. 
 
 Alloys, of bismuth, 181; of 
 tin, description, 245; of tin, 
 list of, 251; of lead, properties, 
 248; of copper, 309; of gold 
 and silver, separation of gold 
 from, 313; gold coins as, 314; 
 of manganese, 321; compo- 
 sition of nickel silver; of 
 nichrome, 331. 
 
 Alum, see potassium aluminium 
 sulfate, 230; use in manufac- 
 ture of matches, 172; proper- 
 ties; preparation, use, 257; as 
 a mordant in dyeing, 257; use; 
 kinds of, 259; chrome, prepa- 
 ration, 317. 
 
 "Alum," for purification of 
 water; see aluminium sulfate, 
 256. 
 
 Aluminium, occurrence, 2.54 ; 
 preparation; properties, use, 
 255; use in welding, 256; occurr- 
 ence ; preparation ; properties ; 
 use ; welding iron, 258 ; manufac- 
 ture by electrolysis; properties; 
 use in Goldschmidt's Thermite 
 Process; in preparation of 
 metals, 278 ; use in preparation 
 of chromium, 316. 
 
 Aluminium hydroxide^ use in 
 purification of water; in dyeing, 
 257. 
 
 Aluminium oxide, use in prepa- 
 ration of ethylene, 199; source 
 of aluminium by electrolysis, 
 255. 
 
 Aluminium sulfate, preparation; 
 "alum;" use, 256; for purifica- 
 tion of water, amount used, 70; 
 use in manufacture of alum, 
 257; use, 259. 
 
 Amalgams, composition; proper- 
 ties, 286. 
 
 Amethyst, form of silicon dioxide, 
 238. 
 
 Ammonia, in soil, 138; from 
 organic matter; formed by aid 
 of catalyzer, 139; properties, 
 solubility, why alkaline, 140; 
 preparation as gas, 141; com- 
 mercial preparation of; syn- 
 thesis of, 142; principles of 
 synthesis of, 142; from action 
 of nitric acid on iron, 145; 
 from organic matter; from 
 manufacture of gas or coke; 
 from direct union of hydrogen 
 and nitrogen, 152; by-product 
 of coking, 189; from calcium 
 cyanide as fertilizer, 266; 
 use in Solvay process, detection 
 in drinking water, 293, 301; 
 use to dissolve silver chloride, 
 311. 
 
 Ammonia-Soda Process, 293. 
 
 Ammonium, why classed as an 
 alkali metal, 299; solubility 
 of salts of, 280; occurrence, 306. 
 
 Ammonium bicarbonate, use in 
 Solvay process, 293. 
 
 Ammonium chloride, prepara- 
 tion, properties, 141, 301, 306. 
 
 Ammonium chloroplatinate 
 preparation, 332. 
 
INDEX 
 
 341 
 
 Ammonium hydroxide, how 
 formed; valence of nitrogen in, 
 141; how formed, 152; prepa- 
 ration, action reversible; use, 
 301; properties, 306. 
 
 Ammonium nitrite, formed by 
 lightning flash, 138. 
 
 Ammonium salts, how prepared; 
 effect of bases on, 141; prepa- 
 ration, 152; list of, 302. 
 
 Ammonium sulfarsenite, prepara- 
 tion ; use in qualitative analysis, 
 179. 
 
 Ammonium sulfantimonite, 181. 
 
 Ammonium sulfate, in fireproof- 
 ing cotton goods, 246. 
 
 Ammonium sulfide, solvent of 
 arsenic trisulfide, 179; solvent 
 of antimony trisulfide, 181. 
 
 Amorphous sulfur, properties, 
 111. 
 
 Amyl acetate, solvent of cellulose 
 nitrates; use of solution, 221. 
 
 Analysis, direct, definition, 3; 
 indirect, 5; groups in, 127; 
 hydrogen sulfide basis of 
 groups in, 112; qualitative, 
 335; quantitative, 337. 
 
 Analysis synthesis, explanation, 
 42, 55. 
 
 Anhydride, definition, 151. 
 
 Aniline dyes, why so called; 
 properties, 233. 
 
 Animal charcoal, preparation ; 
 use, 189. 
 
 Anions, definition, 70; in com- 
 plex cyanides, 217. 
 
 Anode, definition, 44, 55. 
 
 Anthracite coal, how formed; 
 properties, 191. 
 
 Antidotes, vapor of alcohol for 
 chlorine gas; soda-lime and 
 
 sodium thiosulfate for war 
 masks; charcoal for masks, 83; 
 for corrosive sublimate, 287; 
 for snake bite, 322. 
 
 Antimonious acid, properties, 181. 
 
 Antimony, occurrence ; prepara- 
 tion; properties, use, 179, 184; 
 use of in alloys, 180; in chlorine 
 gas, 83; oxides of, 181, 184; 
 reduction of sulfide of, 277. 
 
 Antimony hydroxide, properties, 
 both acid and base, 181, 184. 
 
 Antimony oxychloride, prepara- 
 tion, 181. 
 
 Antimony pentachloride, prepara- 
 tion, 83. 
 
 Antimony trichloride, preparation, 
 83 ; preparation ; properties, 
 180; hydrolysis of, 184. 
 
 Antimony trisulfide, occurrence, 
 preparation, properties; solu- 
 tion in ammonium sulfide. 
 179, 181, 184. 
 
 Antiseptic, mercuric chloride as, 
 287. 
 
 Antitoxins, how formed; use, 232, 
 236. 
 
 Apatite, composition, 170. 
 
 Appolinaris water, how treated, 
 68. 
 
 A'qua regia, solvent of gold or 
 platinum, 146, 152, 314, 332. 
 
 Aqueous solutions, presence of 
 hydrogen ions and hydroxide 
 ions, 72. 
 
 Argon family, properties, 96; in 
 air, how discovered ; properties, 
 156, 160; exception in periodic 
 table, 168. 
 
 Arsine, preparation; properties; 
 test of arsenic poisoning, 177, 
 183. 
 
342 
 
 INDEX 
 
 Arsenic, occurrence, properties, 
 177, 183. 
 
 Arsenic acid, preparation, salts 
 of, 179, 183. 
 
 Arsenic trichloride, preparation; 
 hydrolysis of, 178, 183. 
 
 Arsenic trioxide, formation, prop- 
 erties; use, 177. 
 
 Arsenic trisulfide, preparation; 
 occurrence; use; how dissolved, 
 179, 183. 
 
 Arsenious acid, properties, salts 
 of, 178. 
 
 Artificial silk, from cellulose 
 nitrate, 221. 
 
 Asbestos, a silicate, 239. 
 
 -ate, meaning of, 85. 
 
 Atmospheric pressure, at sea 
 level, 36. 
 
 Atoms, definition, 29; not parts 
 by weight, 52; composite na- 
 ture of, 163. 
 
 Atomic number, how discovered; 
 relation to atomic weight, 168. 
 
 Atomic theory, 50. 
 
 Atomic volume, definition, 169; 
 how determined, 168. 
 
 Atomic weights, definition, 49; 
 told by formula, 52; give pro- 
 portion of substances in com- 
 bination, 56; in table of family 
 groups, 100; classification of 
 families by, 107; and combin- 
 ing volumes, 131; use for com- 
 position of compounds, 132; 
 how determined, 133, 137; 
 relation to periodic system, 
 133; three exceptions in table 
 of periodic system; relation 
 of atomic number to, 168; de- 
 termination of for calcium, 266; 
 table for carbon family, 237. 
 
 Atropine, use, 233. 
 
 Avogadro's law, 135; comparison 
 
 with law of Dulong and 
 
 Petit, 268. 
 
 Babbitt metal, alloy of lead, 248; 
 composition, 180, 245. 
 
 Bacteria, aid in fixing nitrogen; 
 in decay of organic matter, 138; 
 causes of disease, 232; killed 
 by radium, 269; oxidation by, 
 14. 
 
 Baking powders, properties, 229; 
 preparation, 235; composition, 
 cost, 259. 
 
 Baking soda, see sodium bi- 
 carbonate, 229. 
 
 Barium, occurrence, properties, 
 268. 
 
 Barium chloride, properties; 
 use in laboratory, 268, 275. 
 
 Barium peroxide, preparation ; 
 properties; use, 268, 275. 
 
 Barium sulfate, use, 110; in- 
 soluble in water; test of sul- 
 furic acid in solutions, 125; 
 properties, 268. 
 
 Barometers, use of mercury in, 
 286. 
 
 Bases, give hydroxide ions in 
 aqueous solutions, 74; as hy- 
 droxyl compounds, 78; action 
 on ammonium salts, 141, 152. 
 
 Beehive coke ovens, 189. 
 
 Beet sugar, see cane sugar, 222. 
 
 Benzene, occurrence ; preparation ; 
 use, 202, 208; derivative of 
 coal tar, 189. 
 
 Bessemer process for manufac- 
 ture of steel, spiegeleisen in, 
 321; change in manufacture 
 by, 326 ; when invented ; char- 
 
INDEX 
 
 343 
 
 acter of, 333; use for copper, 
 309. 
 
 Bessemer steel, manufacture, 
 326; use, 327. 
 
 Bibasic acids, form acid salts, 76. 
 
 Binary compounds, names of, 15. 
 
 1 'Biscuit" of earthenware. 
 
 Bismuth, occurrence; properties; 
 use; alloys of, 180; compounds 
 of; basic nitrate of, 182, 184. 
 
 Bismuth, basic nitrate, prepara- 
 tion; use, 182, 184. 
 
 Bismuth chloride, properties, 
 hydrolysis of, 182. 
 
 Bismuth nitrate, properties, hy- 
 drolysis, 182. 
 
 Bismuth, " subnitrate," 182. 
 
 Bismuth sulfide, properties, 182. 
 
 Bismuth trioxide, properties ; how 
 dissolved, 182, 184. 
 
 Bituminous coal, properties, 191, 
 192; composition of sample of, 
 161. 
 
 "Black lead," popular name 
 for graphite. 187. 
 
 lack oxide of manganese, use, 
 properties, 321. 
 
 .31ast furnace, production of 
 pig iron, method of operation, 
 324; when invented; use of 
 coal in, 277. 
 
 Bleaching, by chlorine gas, 84, 
 90; by sulfur dioxide; by 
 chlorine, 116; use of hydrogen 
 peroxide, 296. 
 
 Bleaching powder, how prepared ; 
 properties; use, 86; for purifica- 
 tion of water, 70. 
 
 Blowpipe, construction and use, 
 207; oxyhydrogen, why tem- 
 perature is limited, 25; acety- 
 lene, use of, 26. 
 
 Blue vitriol, name for copper 
 
 sulfate, 285; preparation; use, 
 
 310, 314. 
 Bone black, preparation; use, 
 
 189, 193. 
 
 Borates, how formed, 254. 
 Borax, occurrence ; derivation 
 
 properties, 253; use in welding, 
 
 254; use as preservative, 254; 
 
 derivation; use, 258; use, 306. 
 Borax bends, how formed; use 
 
 in laboratory, 254, 258. 
 Bordeaux mixture, composition, 
 
 310. 
 Boric acid, use as preservative; 
 
 as eye-wash ; properties ; prepa- 
 
 ration, 253, 254, 258. 
 Boric anhydride, in forming bor- 
 
 ates, 254. 
 Boron, properties ; occurrence, 
 
 253, 258. 
 
 Boyle, law of, statement ; explana- 
 tion, 36, 40. 
 Brass, alloy of copper and zinc, 
 
 285. 
 
 Bread, manufacture, 229, 235. 
 Brick, 257. 
 
 Brines, where found, 68. 
 British thermal units, equivalent 
 
 in calories, 209. 
 Bromides, occurrence, 98. 
 Bromine, occurrence; prepara- 
 tion; properties; compounds, 
 
 93, 98. 
 
 Bronze, composition, 245. 
 Bunsen burner, flame separated; 
 
 products of flame, 205, 208; 
 
 temperature of flame, 206. 
 Burns, by sulfuric acid, remedies, 
 
 124; from radium, 269. 
 By-products coke ovens, pnxU 
 
 ucts of, 189. 
 
344 
 
 INDEX 
 
 Cadmium, properties ', use in fusi- 
 ble alloys, 285, 288. 
 
 Cadmium hydroxide, properties, 
 260. 
 
 Calcium, occurrence, 260. 
 
 Calcium, preparation; properties; 
 use in laboratory, 261; deter- 
 mination of atomic weight 
 of, 266. 
 
 Calcium aluminate, in manufac- 
 ture of cement, 263. 
 
 Calcium bicarbonate, presence in 
 water; how removed, 264; 
 cause of temporary hardness 
 in water, 275. 
 
 Calcium carbide, preparation ; 
 source of acetylene, 200; prepa- 
 ration, 208; manufacture, use, 
 265; calcium cyanamide from, 
 266. 
 
 Calcium carbonate, from com- 
 bustion, experiment, 11; (lime- 
 stone), effect on water, 62; 
 occurrence, 260; how formed 
 in mortar, 262; deposit from 
 "hard" water, 264; use in 
 Le Blanc soda process, 292; 
 properties, 294. 
 
 Calcium chloride, in apatite, 171 ; 
 preparation; properties; use, 
 264; anhydrous, use, 275; anhy- 
 drous, source of calcium, 261 ; by- 
 product of Solvay process, 293 
 
 Calcium cyanamide, preparation ; 
 use, 266. 
 
 Calcium fluoride, properties, 95; 
 in apatite, 170. 
 
 Calcium hydroxide, for purifica- 
 tion of water, 70; reaction with 
 chlorine, 86; by-product in 
 forming acetylene, 200; prepa- 
 ration, use, 262, 275. 
 
 Calcium light, definition, 26. 
 
 Calcium nitrate, in soil, 138; 
 source of, 139. 
 
 Calcium oxide, manufacture of, 
 261. 
 
 Calcium phosphate, in apatite, 
 170; occurrence, 170, 260; in 
 manufacture of phosphorus, 
 171 ; occurrence ; use ; properties, 
 265; use, 275. 
 
 Calcium silicate, in manufacture 
 of cement, 263. 
 
 Calcium sulfate, use, 110; anhy- 
 drous, properties, 264; cause 
 of permanent hardness in water ; 
 of scale in boilers; how re- 
 moved; properties, 264, 275. 
 
 Calcium sulfide, in Le Blanc soda 
 process, 292. 
 
 Calcium sulfite, preparation, use, 
 117, 118. 
 
 Calculation of amounts of sub- 
 stances in a reaction, 54; of 
 percentage composition from 
 formula, 53. 
 
 Calomel, preparation; use solu- 
 bility, 287. 
 
 Calorie, definition; small calorie, 
 58; equivalent in British ther- 
 mal units, 209. 
 
 Calorimeter, respiration, use of, 13. 
 
 Candle flame, description, 205. 
 
 Cane sugar or sucrose, prepara- 
 tion, 222. 
 
 Cannel coals, properties, use, 192. 
 
 Canon Diablo, diamonds from, 
 186. 
 
 Carbohydrates, definition, 219, 
 234. 
 
 Carbolic acid, preparation; use, 
 228, 235; use of benzene in 
 manufacture of, 202. 
 
INDEX 
 
 345 
 
 Carbon, allotropic forms of; amor- 
 phous, preparation, properties, 
 187, 188, 193. 
 
 Carbon family, occurrence; prop- 
 erties; table of atomic weights 
 of, 237; occurrence; members 
 of, 250. 
 
 Carbon, in manufacture of phos- 
 phorus, 171; position among 
 the elements; properties; oc- 
 currence, 185; gas, preparation; 
 properties, use, 191; how ob- 
 tained by plants; 192; heat 
 of combustion of, 209; relation 
 to living matter, 237; reduc- 
 tion from oxides, 238; use in 
 Le Blanc soda process, 292; 
 use in reduction of oxides, 
 277. 
 
 Carbon dioxide, from combustion 
 experiment, 11; product of 
 oxidation in body, 13; use 
 for effervescent waters, 68; in 
 air, test for, 155; source of, 156; 
 support of vegetation; amount 
 in air, 156; in air, source of; 
 how removed, amount of test 
 of ventilation, 160; in air, not 
 directly harmful, amount per- 
 missible, 161; source of carbon 
 in plants, 192; in illuminating 
 gas, 203; preparation, 212, 216; 
 properties; cause of efferves- 
 cence in soda water, 213; 
 laboratory use of solid form, 
 214; why solution is slightly 
 acid, 213; reactions of; lique- 
 faction of; temperature and 
 vapor pressure of solid form, 
 214; conditions of escape from 
 solution; melting and sub- 
 liming points, 217; in manufac- 
 
 ture of white lead, 250; use 
 in Solvay process, 293. 
 
 Carbon disulfide, preparation; 
 properties; use, 214, 217. 
 
 Carbon monoxide, formation of; 
 preparation, 211; from im- 
 perfect combustion of gasoline, 
 198; in illuminating gas, 203; 
 in water gas, 203; properties; 
 why so poisonous, 212; form- 
 ation of; properties, poison- 
 ous quality, 216; from potas- 
 sium ferri cyanide, 217. 
 
 Carbonic acid, effect on water, 
 62. 
 
 Carborundum, preparation ; prop- . 
 erties; use, 238, 251, 255. 
 
 Casein, occurrence, 230, 235. 
 
 Cassiterite,occurrence ; reduction, 
 241. 
 
 Cast iron, manufacture ; use ; prop- 
 erties, 325. 
 
 Castner-Kellner apparatus, for 
 manufacture of sodium hydrox- 
 ide, 294. 
 
 Catalysis, definition, 9; of de- 
 composition of potassium chlo- 
 rate, 9; use of platinum with 
 sulfur dioxide and oxygen, 119; 
 use of nitrogen dioxide in lead 
 chamber process, 121; osmium 
 and uranium in forming am- 
 monia, 140; why osmium and 
 uranium aid in forming am- 
 monia, 142; use of platinized 
 asbestos in preparation of nitric 
 oxide from ammonia, 143; 
 of decomposition of oxalic acid 
 by sulfuric, 211; enzymes as 
 catalytic agents, 224; use of 
 platinum for, 332. 
 
 Catalyst, definition, 9. 
 
346 
 
 INDEX 
 
 Catalytic agent, manganese diox- 
 ide as, 9; copper chloride as, 
 82; copper oxide as, 307. 
 
 Cathode, definition, 44, 55. 
 
 Cation, definition, 70. 
 
 Cavendish, experimental evi- 
 dence of argon, 156. 
 
 Celluloid, from cellulose nitrates; 
 inflammable, 221. 
 
 Cellulose, occurrence, 234; prop- 
 erties; food for herbivorous 
 animals, 219; fibers containing, 
 220. 
 
 Cellulose hexanitrate : properties ; 
 preparation; use, 220. 
 
 Cellulose nitrate, preparation, 
 220. 
 
 Cement, composition; properties, 
 275; manufacture, 262; com- 
 position, 263; per cents, of 
 materials in; action of, 263; 
 for glass and porcelain, 240. 
 
 Chalcocite, mineral containing 
 copper, 308. 
 
 Chalcopyrite, mineral containing 
 copper, 110, 308. 
 
 Chamber process, see lead cham- 
 ber process, 121. 
 
 Charcoal, burning in oxygen, 10; 
 preparation ; use ; properties, 
 kinds of, use in making iron, 
 188; animal, preparation; use, 
 189, 193. 
 
 Charles, law of, 35, 40. 
 
 Chemical activity, in solutions, 65. 
 
 Chemical nomenclature, binary 
 compounds, 15; oxygen acids 
 of chlorine, 84. 
 
 Chemistry, definition, 2, 5. 
 
 Chile saltpeter, see sodium ni- 
 trate, 291 ; source of nitric acid, 
 143, 152. 
 
 Chlorates, from hypochlorites, 88. 
 
 Chloric acid, from potassium 
 chlorate ; decomposition, 87, 88. 
 
 Chloride, when hydrolyzed, when 
 dissolved and ionized, 180. 
 
 Chlorides, how prepared, 82. 
 
 Chloride of lime, how prepared: 
 properties, 86; composition, 
 275. 
 
 Chlorides of phosphorus, hydro- 
 lysis of, 176. 
 
 Chlorine, antidotes for, 83. 
 
 Chlorine, liquid, for purification 
 of water; amount used, 70; 
 properties, 83, 90; by electro- 
 lysis of sodium chloride, 73, 
 295; oxygen acids of, 84; salts 
 of oxygen acids of, 85; reaction 
 with calcium hydroxide, 86; 
 gas, use in bleaching, 84, 90, 
 116; from chloric acid, 88; 
 for treatment of scrap tin, 246; 
 source of, 291. 
 
 Chlorine dioxide, from chloric 
 acid, 88; use in matches, 89. 
 
 Chloroauric acid, 315. 
 
 Chloroplatmic acid, preparation, 
 332. 
 
 Choke-damp, name for carbon 
 dioxide, 213. 
 
 Cholera, from impure water, 64. 
 
 Chrome alum, preparation, 317, 
 319. 
 
 Chrome green, composition, 319. 
 
 Chrome iron ore, properties, 316. 
 
 Chrome steel, use, 319. 
 
 Chrome tanning, use of potas- 
 sium dichromate in, 317, 318. 
 
 Chrome yellow, properties; use, 
 preparation, 250, 252, 317. 
 
 Chromic anhydride, preparation; 
 properties, 317. 
 
INDEX 
 
 347 
 
 Chromic hydroxide, preparation; 
 properties, 318; presence in 
 chrome leather, 318. 
 
 Chromic oxide, in preparation of 
 chromium, 278, 316; prepara- 
 tion; properties, 318. 
 
 Chromite, see chrome iron ore, 
 316. 
 
 Chromium, occurrence, prepara- 
 tion, properties, use, 316; prepa- 
 ration, 278; use in nichrome, 
 331. 
 
 Cinnabar, mineral source of mer- 
 cury, 285. 
 
 Citric acid, in lemons, 20. 
 
 Clay, composition of, formation, 
 255; when plastic; use, 241; 
 composition; fusion; products 
 of; manufacture of brick, earth- 
 'enware, porcelain, 257; use in 
 manufacture of brick, earthen- 
 ware, porcelain, 239. 
 
 Clay iron stone, occurrence, 324. 
 
 Coal, how formed; varieties; 
 properties, 191 ; cannel, coking 
 and non-coking, 192; bitu- 
 minous, properties; anthracite, 
 how formed, properties, 191; 
 table of changes in, 192; bitu- 
 minous, kinds, 193. 
 
 Coal tar, from coking retorts, 
 products derived from, 189; 
 products of, 208. 
 
 Coal tar dyes, properties, 234. 
 
 Cobalt, exception in periodic 
 table, 168; properties; com- 
 pounds of, 331. 
 
 Cobalt chloride, properties; use 
 331. 
 
 Cobalt oxide, properties; use, 
 331. 
 
 Cocaine, use, 233. 
 
 Coke, preparation; use, 189, 193. 
 preparation; use, 189; use 
 in blast furnace, 324. 
 
 Coke ovens, beehive, 189; to 
 recover by-products, 189. 
 
 Coking coals, properties, use, 
 192. 
 
 Collection of gases, 9. 
 
 Collision between elastic bodies, 
 law of reaction, 32. 
 
 Collodion, solution of nitrate of 
 cellulose, 221. 
 
 Colloidal solutions, description; 
 preparation, 240; description; 
 preparation, 251. 
 
 Colloids, description; separation 
 in digestion; effect upon clay, 
 241. 
 
 Combining proportions, law of, 
 48, 55, 132. 
 
 Combining volumes and atomic 
 weights, 131. 
 
 Combustion, explanation of, 6; 
 quantity of heat of, 13; relation 
 of air to, 6; spontaneous, how 
 caused: how avoided, 14; effi- 
 ciency in boilers, 194; table of 
 heats of combustion for illumi- 
 nating gas, 208. 
 
 Combustion, heat of, definition. 
 17; for acetylene compared 
 with carbon and hydrogen, 201 ; 
 for carbon and methane, 210. 
 
 Complex cyanides, 215. 
 
 Compounds, definition, 4, 5; 
 in terms of atomic weights, 
 132; formation of volatile, in- 
 soluble and slightly ionized, 
 279. 
 
 Concentrated sulfuric acid, cau- 
 tion in use of, 124; use in 
 preparation of ethylene, 199. 
 
-348 
 
 INDEX 
 
 Conservation of energy, law of, 14. 
 
 Constant proportion, law of, 
 3, 48, 55. 
 
 "Contact process," in manufac- 
 ture of sulfur trioxide and sul- 
 furic acid, 119. . 
 
 Convection, in water, 60. 
 
 Copper, use in preparation of 
 sulfur dioxide, 116; specific 
 gravity, 169; alloys of, 180; 
 reduction of sulfide of; refining 
 by electrolysis, 278; occurrence; 
 preparation, minerals contain- 
 ing copper, 308; impurities 
 in; electrolytic reduction; prop- 
 erties; use; conductivity, how 
 impaired, salts of, 309; per- 
 centage in U. S. silver coins, 
 311; properties; compounds of r 
 preparation, 313; percentage 
 in nickel coins, 331; use in 
 nickel silver, 331. 
 
 Copper acetate, in Paris green, 
 179. 
 
 Copper arsenite, in Paris green, 
 179. 
 
 Copper chloride, as catalytic 
 agent, 82. 
 
 Copper Matte, mixture of sulfides 
 from sulfide ores; properties; 
 reduction, 309. 
 
 Copper oxide, reduction, 24, 26; 
 use in determining composition 
 of water, 45; in preparation of 
 sulfur dioxide, 116; as catalytic 
 agent, 307. 
 
 Copper pyrites, see chalcopyrite 
 308. 
 
 Copper sulfate, by-product in 
 preparation of sulfur dioxide, 
 116; popular name of, 285; 
 preparation; use, 310. 
 
 Copperas, preparation, 329, 333. 
 
 Corrosive sublimate, solubility; 
 antidote; use, 287, 288; prepa- 
 ration, 289. 
 
 Corundum, occurrence; use, 255. 
 
 Cream of tartar, see acid potas- 
 sium tartrate, 229. 
 
 "Creosote" for treating lumber, 
 189. 
 
 Cryolite, in electrolysis of alumin- 
 ium oxide, 255. 
 
 Crystallization, purification by, 2. 
 
 Crystalloids, description, 241. 
 
 Cubic centimeter, defined, 28. 
 
 Cupric chloride, preparation, 83. 
 
 Cupric oxide, in preparation of 
 sulfur dioxide, 116. 
 
 Cuprous chloride, preparation, 83; 
 properties, 310; preparation; 
 solubility, 314. 
 
 Cuprous oxide, preparation; test 
 for glucose, 309, 314. 
 
 Cuprous sulfide,. in copper 
 matte, 309. 
 
 Cyanides, poisonous quality, 217; 
 soluble as complex cyanides, 
 215. 
 
 Cyanide process, in gold mining, 
 312. 
 
 Davy safety lamp, construction 
 of, why safe, 197; why in- 
 vented, 208. 
 
 Davy, Sir Humphrey, preparation 
 of potassium and sodium by 
 electrolysis, 278. 
 
 Deliquescence, definition, 69. 
 
 Denatured alcohol, preparation; 
 use, 226. 
 
 Developer in photography, 312. 
 
 Dextrin, preparation; use, 224, 
 234. 
 
INDEX 
 
 349 
 
 Di-, prefix, 174. 
 
 Diabetes, change of sugar in, 223. 
 
 Dialysis, definition, 241; use. 
 251. 
 
 Diamonds,occurrence; properties; 
 how formed, 185, 192; arti- 
 ficial, 186; use, 187. 
 
 Diastase, how formed ; properties, 
 224. 
 
 Dibasic acids, definition, 76, 118. 
 
 Diffusion of gases, experiment, 
 29; explained by kinetic theory, 
 31. 
 
 Direct analysis, definition, 3, 5. 
 
 Disease germs in water, 64. 
 
 Disinfectants, sulfur dioxide; 
 formaldehyde, 115; zinc chlo- 
 ride, 288; potassium perman- 
 ganate as r 322. 
 
 Disodium phosphate, ionization 
 of, 176; product in digestion; 
 function in body, 291. 
 
 Dissociation, defined, 150, 153; 
 of water, 25; of nitrogen 
 tetroxide, 150. 
 
 Distillation, purification by, 2. 
 
 Divisions of Group I, contrasts 
 between elements of, 308. 
 
 Divisions of Group VI, contrasts 
 between elements of, 316, 318. 
 
 Dolomite, composition, 261; com- 
 position, 283. 
 
 Drummond light, definition, 26. 
 
 "Dry plates," preparation; use 
 311. 
 
 Dulong and Petit, law of, for 
 exceptional elements, 137; law 
 of, comparison with law of 
 Avogadro, 267 ; table for appli- 
 cation of, 268. 
 
 Dust, explosive mixtures with air, 
 198. 
 
 Dyeing, potassium dichromate 
 as mordant in, 317. 
 
 Dyes, natural, source, list of; 
 artificial, aniline; properties, 
 233; use of hydrocarbons in 
 manufacture of, 202; vege- 
 table and artificial, 236. 
 
 Dynamite, from nitroglycerine, 
 125; manufacture, how ex- 
 ploded, 228, 235. 
 
 Earthenware, 257. 
 
 Efflorescence, definition, 69. 
 
 Eggs, preservation, use of sodium 
 silicate, 240, 296. 
 
 Elastic bodies, law of rebound, 
 32, 40. 
 
 Electric arc, temperature of, 25; 
 use in manufacture of nitric 
 oxide and nitric acid, 149. 
 
 Electric current, between metals, 
 cause of, 242; why direction of 
 current is wrongly stated, 245. 
 
 Electrodes, definition, 43, 55; gas 
 carbon for, 191. 
 
 Electrolysis, definition, 43; of 
 sulfuric acid, 42, 55; explana- 
 tion, 65; of sodium chloride, 
 fused; in aqueous solution, 
 72, 78; in manufacture of 
 aluminium, 255 ; in preparation 
 of calcium, 261; of salt solu- 
 tion, Castner-Kellner apparatus 
 for, 294; of salt solution, 305. 
 
 Electrolyte, definition, 44; separa- 
 tion into ions, 65. 
 
 Electrolytic methods, for refining 
 copper; for manufacture of 
 aluminium, sodium and mag- 
 nesium, 278. 
 
 Electromotive series, definition, 
 251; explanation; table of^ 243;. 
 
350 
 
 INDEX 
 
 to determine position of metals 
 in, 244. 
 
 Electrons, definition; constituent 
 part of atoms of some elements; 
 weight of, 163; relation to 
 potential between metals; to 
 electric current between metals, 
 242; in radium disintegration, 
 272. 
 
 Electropositive, definition, 243. 
 
 Electroscope, gold leaf, to detect 
 radium, 269. 
 
 Elements, definition, 4; example, 
 number of, 4, 5; most com- 
 mon, 4; groups of, 92; families 
 of, properties, table, 100; rate 
 of combination of, dependent 
 on temperature, 120; disinte- 
 gration of, 270; how detected 
 in stars, 303. 
 
 Emery, see corundum, 255. 
 
 Endothermic compounds, reac- 
 tions, definition, 201. 
 
 Energy, kinetic; potential, 2; 
 law of conservation of, 14; 
 indestructibility, 17; produc- 
 tion of, from food, 230; from 
 radium, 269. 
 
 Enzymes, description, 224. 
 
 Equations, explanation ; writing 
 of, 52; what they represent, 
 56; how combined, 86. 
 
 Equilibrium, between water and 
 vapor, 69; in reactions, 74; 
 in preparation of sulfur triox- 
 ide, 118; in reversible reaction, 
 effect of temperature on, 
 121. 
 
 Etching glass, by hydrofluoric 
 acid, 96. 
 
 Ethyl alcohol, preparation, 224, 
 235. 
 
 Ethylene, preparation ; occur- 
 rence; compounds; why termed 
 unsaturated ; gives luminous 
 quality to gas, 199, 208. 
 
 Ethylene bromide, preparation, 
 199. 
 
 Ethylene chloride, preparation, 
 199. 
 
 Explosive mixtures, composition, 
 197. 
 
 Explosives, dependence on nitric 
 acid, 143, 152. 
 
 False gold leaf, composition, 83. 
 
 Fats, substances composed of; 
 acids in; salts derived from, 
 227; acids in; saponification, 
 235. 
 
 Fatty acids, salts of, 227. 
 
 Ferric chloride, preparation; prop- 
 erties, 330; anhydrous, prepa- 
 ration, 333. 
 
 Ferric hydroxide, preparation, 
 330. 
 
 Ferric oxide, use in welding, 256; 
 use in Goldschmidt's Thermite 
 process, 278. 
 
 Ferric sulfate, preparation, 330, 
 333. 
 
 Ferrocyanogen ion, description, 
 216. 
 
 Ferro-manganese, composition 
 use; preparation, 321. 
 
 Ferrosilicon, manufacture; prop- 
 erties, 238. 
 
 Ferrous chloride, how formed, 21. 
 
 Ferrous cyanide, in preparation 
 of potassium ferrocyanide, 215. 
 
 Ferrous hydroxide, preparation, 
 properties, 329, 333. 
 
 Ferrous silicate, in reduction of 
 copper. 309. 
 
INDEX 
 
 351 
 
 Ferrous sulfate, how formed, 21; 
 electrolysis of, 216; popular 
 name of, green vitriol, 285; 
 preparation, 329. 
 
 Ferrous sulfide, precipitation of, 
 115; hydrogen sulfide from, 112; 
 in copper matte, 309. 
 
 Fertilizers, from mineral phos- 
 phates, use of sulfuric acid, 125; 
 use of gypsum, 261; forms of 
 phosphates used as, 265; use 
 of ammonia from calcium 
 cyanide, 266; forms of phos- 
 phates used as, 275; use of 
 sodium nitrate as, 291; potash 
 salts in, 298. 
 
 Fire, 6; early methods of get- 
 ting, 172. 
 
 Firedamp, popular name for 
 methane; cause of explosions 
 in mines, 197. 
 
 Fireproofing fabrics, with soluble 
 glass, ^40; of cotton goods with 
 sodium stannate, 246; use of 
 sodium silicate, 296. 
 
 Flames, what luminosity de- 
 pends upon, 205; temperature 
 of, 206; blowpipe, oxid zing, 
 reducing, 207, 209; what lum- 
 inosity depends upon, 208; 
 colored, how produced, 302, 
 306. 
 
 Flash light powders, composition, 
 283. 
 
 Flashing point, of gasoline, of 
 kerosene, 208. 
 
 Flint glass, composition, 239, 
 
 Flowers of sulfur, definition, 110. 
 
 Fluorine, occurrence, properties; 
 how prepared, 95, 98; prepara- 
 tion, 99. 
 Fluorite, use as flux, 261. 
 
 Fluorspar, properties, 95; source 
 of fluorine; use as flux, 201. 
 
 Food, production of heat and 
 energy from, 230, 235; for 
 plants, sources of, substances, 
 231 ; formation of tissues from, 
 231, 235. 
 
 Formulas for molecules, 51 ; what 
 they designate, 56; graphical, 
 103; represent gram molecular 
 volume, 135; of elementary 
 gases, how determined, 136; 
 for nitrogen tetroxide and 
 nitrogen dioxide at different 
 temperatures, 151 ; for va- 
 lences in different groups, 
 166. 
 
 Frasch, Hermann, process for 
 getting sulfur in Louisiana, 109. 
 
 Fraunhofer lines, discovery of, 
 explanation of, 303, 307. 
 
 Fructose, how prepared; prop- 
 erties, 223. 
 
 Furnace, Bessemer, 327; open 
 hearth, 328; reverberatory, use 
 in manufacture of wrought 
 iron, 326. 
 
 Fusible alloys, use of cadmium in, 
 288. 
 
 Galena, properties, 247; reduc- 
 tion of, 310. 
 
 Galvanized iron, how made; when 
 poisonous, 284. 
 
 Garnet, a silicate, 239. 
 
 Gas, determination of weight of, 
 38, 40; measure of pressure of, 
 40; law of partial pressures, 
 61 ; illuminating, properties, 
 203; water, properties, 203; 
 producer, manufacture, use, 
 204. 
 
352 
 
 INDEX 
 
 Gases, collection and storage of, 
 9; illustration of diffusion, 29; 
 kinetic theory of, 31; diffusion 
 explained by kinetic theory, 31 ; 
 law of change of volume with 
 temperature, 35; law of change 
 of volume with pressure, 37; 
 reduction to standard volume, 
 37; diffusion explained by 
 kinetic theory, 40; combina- 
 tion by volume, diagram, 130; 
 combination by volume, Gay 
 Lussac's law, 131, 136; "per- 
 fect," 132; why "noble," so- 
 called, 158. 
 
 Gas Carbon, preparation ; proper- 
 ties, use, 191, 193. 
 
 Gasoline, composition; use, 198. 
 
 Gasometer, method of use, 9. 
 
 Gastric juice, dependence on 
 salt, 291. 
 
 Gay Lussac's law of combination 
 by volume, 131. 
 
 Gay-Lussac tower, use in manu- 
 facture of sulfuric acid, 123. 
 
 Germicide, acid sodium sulfite 
 as, 118; use of sulfur dioxide 
 as, 115; use of potassium per- 
 manganate, 322. 
 
 Germs, carried by water, how 
 removed, 64, 69. 
 
 Glass, composition; manu- 
 facture; properties; use; 
 varieties, 239, 251; etching by 
 hydrofluoric acid, 96. 
 
 Glazes, for clay products, 258, 
 259. 
 
 Glover tower, use in manufacture 
 of sulfuric acid, 123. 
 
 Glucose, how prepared; use; in 
 diabetes, 223, 234; use in pre- 
 paring cuprous oxide- 309. 
 
 Gluten, preparation, 230, 235. 
 
 Glycerine, see glycerol, 227. 
 
 Glycerol, from fats, 228, 235; 
 manufacture; use, 227, 235. 
 
 Glyceryl, in fats, 227, 235. 
 
 Gold, action of aqua regia upon, 
 146; occurrence, separation, 
 312, 314; percentage in coins, 
 properties, oxides of; chlorides 
 of, 313. 
 
 Goldschmidt's thermite process, 
 use, 256. 
 
 Gram, defined, 28. 
 
 Gram-atom, how designated, 56; 
 simple numbers in unit volume, 
 132; weight of, 137. 
 
 Gram-molecule, how designated, 
 56; weight of, volume of, 134. 
 
 Gram molecular volume, meas- 
 ure of, 134, 136. 
 
 Granite, composition, 239. 
 
 Graphical formulas, 103; to ex- 
 press valence, 107. 
 
 Graphite, occurrence, prepara- 
 tion; use, properties, 187, 193. 
 
 Gravitation, law of, 1. 
 
 Green vitriol, see ferrous sulfate, 
 285. 
 
 Group zero, 155. 
 
 Group I, first division, 290. 
 
 Group I, second division, 308. 
 
 Group II, first division, 260. 
 
 Group II, second division, 283. 
 
 Group III, 253. 
 
 Group IV, 185, 237. 
 
 Group V, 138, 170. 
 
 Group VI, 109. 
 
 Group VI, second division, 316. 
 
 Group VII, 92. 
 
 Group VII, second division, 320. 
 
 Group VIII, characteristics of; 
 valence, 323. 
 
INDEX 
 
 353 
 
 Groups of elements, 92 ; of metals 
 in qualitative analysis, 112, 
 127; of periodic system, 162. 
 
 Gun cotton, manufacture by sul- 
 fur ic acid, 125; preparation; 
 properties, 220. 
 
 Gunpowder, composition, cause 
 of explosion, how rate of burn- 
 ing is regulated, 300, 306; 
 preparation of saltpeter for, 279. 
 
 Gypsum, use, 110; properties, 
 use, 261; in manufacture of 
 plaster of Paris, 264. 
 
 Half-life period of radioactive 
 
 elements, 276. 
 Half-metals, distinguished from 
 
 non-metals and metals, 237. 
 Halogen family, list; atomic 
 
 weights, 92, 98; properties of 
 
 compound of, 97, 99. 
 <: Hard finish," use of plaster of 
 
 Paris in, 264. 
 
 "Hard" water, explanation, 263. 
 Heat of combustion, definition, 
 
 13, 17. 
 Heat, production of, from food, 
 
 230. 
 Helium, boiling point, 35; where 
 
 found; how formed, 158, 160; 
 
 constituent part of atoms of 
 
 some elements; weight of, 
 
 163; by decomposition of 
 
 radium; cause of radioactive 
 
 phenomena, 270. 
 Hematite, ore of iron, 324. 
 Henrys law of gas pressure, 61; 
 
 of absorption of gas by water, 
 
 159; applied to carbon dioxide, 
 
 213. 
 High-speed tools, tungsten in, 
 
 318. 
 
 23 
 
 Hunyadi water, mineral content 
 of, 283. 
 
 Hydrates, definition, 68. 
 
 Hydraulic cement, see cement, 
 263. 
 
 Hydraulic mining of gold, 312. 
 
 Hydriodic acid, use in prepara- 
 tion of phosphonium iodide, 
 173. 
 
 Hydrocarbons, number of; marsh 
 gas series of, 195; structure 
 of methane series of; graphic 
 formulas, 196; list of series and 
 compounds, 202; valences of 
 carbon and hydrogen in, 207. 
 
 Hydrochloric acid, with zinc; 
 with iron, 21; how prepared; 
 properties, 80, 90; boiling point 
 of solutions, 81, 90; action of 
 oxidizing agents on, 82; action 
 on metals, 81; action on 
 hydroxides and oxides, 82, 90; 
 standard for unit volume, 132; 
 by hydrolysis of chlorides of 
 phosphorus, 176; source in 
 .digestion, 291. 
 
 Hydrocyanic acid, preparation ; 
 use; poisonous quality, 215, 
 217. 
 
 Hydroferrocyanic acid, relation 
 to potassium ferrocyanide; 
 preparation of hydrocyanic 
 acid from, 215. 
 
 Hydrofluoric acid, how prepared; 
 use in etching glass, 96, 99. 
 
 Hydrogen, by passing steam over 
 red hot iron; by action of 
 sodium or potassium on water, 
 19; in substances of acid taste, 
 20; from acids and metals, 21; 
 danger of explosion; proper- 
 ties, 21; properties; explosive 
 
354 
 
 INDEX 
 
 mixtures; occurrence; reduc- 
 tion by, 22; liquid, solid, 23; 
 by action of sodium or potas- 
 sium on water, 26; properties; 
 occurrence, 26; velocity of 
 molecules, 33; diffusion ex- 
 plained by kinetic theory, 33; 
 as illustration of absolute 
 temperatures, 34; liquefying 
 temperature, 34; weight of liter 
 of, 43; ions, in aqueous solu- 
 tions, 72; ions, from acids, 73; 
 replaced by metals, 73, 78; 
 weight of liter of, 99; valence 
 of, 105. 
 
 Hydrogen ions, from nitric acid, 
 114- 
 
 Hydrogen peroxide, properties, 
 49; preparation from barium 
 peroxide, 275; preparation, use 
 in bleaching, 296. 
 
 Hydrogen sulfide, preparation, 
 properties; occurrence; use in 
 chemical analysis, 112, 127; 
 occurrence; use in chemical 
 analysis, 127; weight of liter, 
 128; in illuminating gas, 203; 
 in preparation of zinc sulfide, 
 285. 
 
 Hydrolysis, definition, 176, 183; 
 of salts of phosphoric acid, 175; 
 of chlorides of phosphorus, 176; 
 of arsenic trichloride, 178; 
 of bismuth compounds, 182; 
 of salts of phosphoric acid, 183; 
 of salt in digestion, 291. 
 
 Hydrosulfides, preparation, 115, 
 127. 
 
 Hydroxide ions, in aqueous solu- 
 tion, 72; from bases, 74. 
 
 Hydroxyl compounds, how ion- 
 ized, 78. 
 
 Hypo-, meaning of, 85. 
 
 Hypochlorites, preparation, 86. 
 
 Hypochlorous acid, formed in 
 bleaching, 84. 
 
 Hypophosphorous acid, mono- 
 basic, 176, 183. 
 
 "Hyposulfite," sodium, 126. 
 
 -ic, meaning of, 84. 
 
 Ice, latent heat of fusion of, 58, 69. 
 
 -ide, ending for binary com- 
 pounds, 15. 
 
 Illuminating gas, manufacture ; 
 composition; 202, 208; what il- 
 luminating quality depends 
 upon, 203; table of heats of 
 combustion for, 209; composi- 
 tion of different kinds, 210. 
 
 Indestructibility of matter, law 
 of, 12. 
 
 India rubber, use of sulfur, 110. 
 
 Indicators, definition ; use, 75, 79. 
 
 Indigo, artificial product, 233. 
 
 Indirect analysis, 5. 
 
 Insoluble compound, conditions 
 for formation of, 279. 
 
 Invert sugar, how prepared, 223, 
 234 ; why so named ; occurrence, 
 223. 
 
 Iodine, occurrence; preparation; 
 properties, 94, 98; tincture of, 
 preparation; use, 94, 98; com- 
 pounds; properties, 98; in thy- 
 roid gland, 95, 98; test for 
 starch, 222. 
 
 lonization, in solutions, 65; of 
 acids in aqueous solutions, 73; 
 of water, 74;in dilute solutions, 
 77; of nitric acid, 114; of 
 phosphoric acid, 175; of diso- 
 dium phosphate, 176; of air by 
 radium, 269. 
 
INDEX 
 
 355 
 
 Ions, in solutions, 65; positive- 
 negative; migration of, 66; 
 in formation of precipitates; 
 formation reversible, 113;ferro- 
 cyanogen, 216; transfer from 
 metals to solution, 242. 
 
 Iron, burning in oxygen, 11; used 
 to decompose water, 19; his- 
 tory of use; blast furnace; im- 
 portance of, 277; reduction of 
 oxides of, 277; importance; 
 history, 323, 333; occurrence; 
 manufacture; ores of, 324; 
 pig or cast, manufacture, use, 
 324; wrought, manufacture, 
 use, 326; salts of, 329. 
 
 Iron carbide, in manufacture of 
 steel, 326. 
 
 Iron pyrites, source of sulfur 
 dioxide, 122; in manufacture of 
 sulfuric acid, 110. 
 
 Isotopes, compounds of, excep- 
 tion to law of constant pro- 
 portion, 3. 
 
 -ite, meaning of, 85. 
 
 Jasper, form of silicon dioxide, 
 
 238. 
 Jelly, manufacture, 224. 
 
 Kaolin, a silicate, 239; how 
 formed-, 255. 
 
 Kerosene, use, flashing point of, 
 199. 
 
 Kilogram, defined, 28. 
 
 Kimberly, diamonds from, 185. 
 
 Kindling temperature, explana- 
 tion, 12, 17. 
 
 Kinetic energy, 2. 
 
 Kinetic theory of gases, explana- 
 tion, 31, 40. 
 
 Kipp generator, use of, 21. 
 Krypton, where found, 158. 
 
 Lacquers from cellulose nitrate, 
 221. 
 
 Lampblack, preparation; use; 
 properties, 188, 193. 
 
 Latent heat of fusion of ice. 58, 
 69; of vaporization of water, 
 59, 69. 
 
 Lavoisier's experiment, propor- 
 tion of oxygen to nitrogen in 
 air, 7. 
 
 Law of gravitation, 1; of con- 
 stant proportion, 3; definition, 
 5; of indestructibility of mat- 
 ter, 12; of conservation of 
 energy, 14; of velocity of 
 molecules in gases, 33; of 
 change in volume of gas 
 with temperature, 35; of 
 Charles, 35, 40. 
 
 Law of Boyle, statement; ex- 
 planation, 36, 40; of change of 
 volume of gas with pressure, 
 37; of multiple proportions, 
 49, 56; of combining propor- 
 tions, 48, 55, 132; of constant 
 proportion, 48, 55; partial 
 pressures of gases (Henry's 
 law), 61; of vapor pressures, 
 63. 
 
 Law, Gay Lussac's for combina- 
 tion of gases, 131; Henry's, 
 of gas pressure, 61; of van't 
 Hoff-Le Chatelier, 120, 142, 
 148; of Avogadro, 135; of 
 Dulong and Petit, for excep- 
 tional elements, 137; of Henry, 
 159; of Dulong and Petit, 
 267; table for application of, 
 268. 
 
356 
 
 INDEX 
 
 Leaching, process described, 297. 
 
 Lead chamber process for manu- 
 facturing sulfuric acid, 121, 
 128, 247. 
 
 Lead, use of alloys of, 180; 
 danger of use for water pipes, 
 247; oc urrence; preparation; 
 use; oxides of, salts of, 247, 252; 
 differences in atomic weight, 
 273; reduction of sulfide of, 
 278; use, 247; alloys of; oxides 
 of, 248. 
 
 Leadacetate,preparation;use,250. 
 
 " Lead burning," 122. 
 
 Lead chloride, solubility; prepa- 
 ration, 250. 
 
 Lead c h r ornate , preparation ; 
 properties; use, 317; see chrome 
 yellow, 252; use, 319. 
 
 Lead dioxide, preparation; use 
 in storage batteries, 248. 
 
 Lead nitrate, preparation, 248, 
 249. 
 
 Lead oxide, preparation, 247. 
 
 "Lead" pencils, made from 
 graphite, 187. 
 
 Lead plumbate, see red lead, 252. 
 
 Lead sulfate, preparation, 247; 
 in storage batteries, 249. 
 
 Lead sulfide, precipitation of, 
 114. 
 
 Le Blanc soda process, three 
 operations of , 291, 305. 
 
 Legumes, fixation of nitrogen by, 
 138, 152. 
 
 Lemons, citric acid in, 20. 
 
 Light, ultra-violet for purifica- 
 tion of water, 70. 
 
 Lignite, properties, 191, 192. 
 
 Lime, manufacture of, 261, 275; 
 "slaking" of, 11; for purifica- 
 tion of water, 70. 
 
 Lime light, definition, 26. 
 
 " Lime -nitrogen, " preparation ; 
 use, 266. . 
 
 Lime -sulfur wash, in spray for 
 vegetation, 110. 
 
 Lime water, preparation, 11. 
 
 Limestone (calcium carbonate); 
 effect on water, 62; composi- 
 tion, 260; use in blast furnace, 
 324. 
 
 Limonite, ore of iron, 324. 
 
 Liquid air, how obtained, 10. 
 
 Liquid hydrogen, properties, 22. 
 
 Liter, defined, 28. 
 
 Liter of a gas, determination of 
 weight of, 38. 
 
 Litharge , preparation , 248 . 
 
 Lithopone, composition, 289. 
 
 Litmus, affected by acids, 11; 
 as an indicator, 76; action 
 of sodium hydroxide on, 20. 
 
 "Liver of sulfur," use by Cav- 
 endish to absorb oxygen, 157. 
 
 Lodestone, properties, 330. 
 
 Lubricant, graphite, 187. 
 
 Lubricating oils, product from 
 petroleum, 199. 
 
 Lunar caustic, see silver nitrate, 
 311. 
 
 Lye, how obtained ; use in making 
 soft soap, 297. 
 
 Magnesite, composition, 283, 330; 
 
 properties, 316; ore of iron, 
 
 324. 
 Magnesium, manufacture by 
 
 electrolysis, 278; occurrence; 
 
 properties of compounds of; 
 
 use; compounds of, 283, 284, 
 
 288. 
 Magnesium ammonium arsenate, 
 
 179. 
 
INDEX 
 
 357 
 
 Magnesium carbonate, occur- 
 rence, 261 ; in mineral form, 283. 
 
 Magnesium sulfate, occurrence; 
 properties, 283; use in 
 manufacture of matches, 172. 
 
 Magnetic oxide of iron, from 
 combustion experiment; prop- 
 erties, 11; of iron, by passing 
 steam over iron, 19; reduc- 
 tion of, 24, 27; preparation, 
 properties, 330; preparation, 
 333. 
 
 Malachite, mineral containing 
 copper, 308. 
 
 Malt, agent in fermentation, 224. 
 
 Maltose, preparation ; properties, 
 223, 234. 
 
 Manganese, occurrence; prop- 
 erties; use, 320, 322; alloys 
 of, 321; oxides of, 321. 
 
 Manganese chloride, preparation, 
 83. 
 
 Manganese dioxide, catalytic 
 agent, 9; to oxidize hydro- 
 chloric acid, 83; occurrence, 
 use, 320, 322. 
 
 Maple sugar, description, 222. 
 
 Marble, how formed, 261. 
 
 Marsh gas, popular name for 
 methane, 196. 
 
 Marsh gas series of hydrocar- 
 bons, 195. 
 
 Matches, use of phosphorus, 172. 
 
 Matter, law of indestructibility, 
 12. 
 
 Mauve, source of, 233. 
 
 Meerschaum, a silicate, 239. 
 
 Meker burner, temperature of 
 flame, 207. 
 
 Melting points of elements, rela- 
 tion to position in periodic 
 system, 168. 
 
 Mercuric chloride, source of 
 calomel; properties; antidote, 
 287. 
 
 Mercuric fulminate, use; prop- 
 erties, 287, 288. 
 
 Mercuric oxide, for preparation 
 of oxygen, 7 ; use in laboratories, 
 286. 
 
 Mercurous chloride, preparation, 
 use, 287, 288. 
 
 Mercury, occurrence ; prepa- 
 ration, properties; use, amal- 
 gams; valence, 285, 286, 288; 
 use in Castner-Kellner appara- 
 tus, 295. 
 
 Mesothorium, use in radiolite 
 watches, 273, 276. 
 
 Metallic properties, variation in 
 periods and groups, 166. 
 
 Metallurgy, definition, 277. 
 
 Metals, replace hydrogen in 
 acids, 73, 78; groups of in 
 qualitative analysis, 112; prop- 
 erties of salts of alkali, 217; 
 solution pressure of, explana- 
 tion; cause of electric cur- 
 rent between metals, 242; 
 reduction of ores of, 277; of 
 second division, group II, 
 characteristics of, 283, 287. 
 
 Metaphosphoric acid, 174. 
 
 Metastannic acid, preparation, 
 145; preparation; use as test 
 in analysis, 246, 251. 
 
 Meter, defined, 28. 
 
 Methane, properties, 196; prepa- 
 ration, 197; occurrence, prepa- 
 ration; explosive mixtures 
 of, 208; heat of combustion of, 
 210. 
 
 Methane series of hydrocarbons, 
 structure of, 196. 
 
358 
 
 INDEX 
 
 Metric weights and measures; 
 
 equivalents, 28, 39. 
 Migration of ions, 66. 
 Milk of lime, preparation, 262. 
 Millimeter, defined, 28. 
 Mineral phosphates, use for 
 
 fertilizers, 125. 
 Mineral waters, natural and 
 
 artificial, 67; mineral content 
 
 of, 283. 
 
 Mispickel, source of arsenic, 177. 
 Mixtures, separation of, 2. 
 Moissan, method of preparing 
 
 fluorine, 96; preparation of 
 
 diamonds, 186. 
 Molecular theory, explanation, 
 
 29, 40. 
 Molecular weight of compound, 
 
 how determined, 134. 
 Molecules, definition, 29, 39; 
 
 law of .velocity in gases, 33; 
 
 number in cubic centimeters, 
 
 33, 40, 273, 276; number in 
 
 equal volume, 135; number in 
 
 unit volume, 135. 
 Mono-, prefix, 174. 
 Monobasic acid, definition, 76, 
 
 118. 
 Morphine, source; use; danger 
 
 of use, 233. 
 Mortar, preparation; use; cause 
 
 of strength, 262. 
 Multiple proportions, law of, 
 
 49, 56. 
 
 Names of binary compounds, 15; 
 
 of oxygen acids of chlorine 
 
 and salts, 84, 85. 
 Naphthalene, derivative of coal 
 
 tar, 189. 
 
 Natural gas, occurrence, 196. 
 Natural waters, 61. 
 
 "Negative" in photography, ex- 
 planation, 312. 
 
 Negative and positive valences, 
 104, 106. 
 
 Neon, where found, 158. 
 
 Neutral solution, definition, 75, 
 79 
 
 Neutralization, definition, 78. 
 
 Nichrome, composition of alloy; 
 use, 331, 333; high melting 
 point, 333. 
 
 Nickel, properties; use; alloy of, 
 330; percentage in coins, 331. 
 
 "Nickel plate," how deposited; 
 use, 330. 
 
 Nickel silver, composition of 
 alloy; use, 331, 333. 
 
 Nicotine, occurrence; properties, 
 233. 
 
 Niton, how formed; properties, 
 158, 161; how produced from 
 radium, 270; properties, 272. 
 
 Nitrates, from decay of organic 
 matter, 144. 
 
 Nitrates of cellulose, uses in 
 solution, 221. 
 
 Nitric acid, ionization, 114; use 
 in manufacture of sulfuric 
 acid, J22, 128; sources of, 143; 
 oxidation with, 144; prepara- 
 tion, properties, 144; action 
 on metals, 145; preparation, 
 use, as oxidizing agent, 152; 
 source of; use, 291. 
 
 Nitric oxide, formed by electricity, 
 139; preparation; properties; 
 use, 148; dependence of forma- 
 tion on temperature, 148; 
 formed from ammonia by 
 use of platinized asbestos 
 as catalyzer, 143, 152; proper- 
 ties; manufacture by means 
 
INDEX 
 
 359 
 
 of electric arc, 148; prepara- 
 tion, 153; presence after dyna- 
 mite explosion, 228. 
 
 Nitrites, decomposition by acids, 
 147, 152. 
 
 Nitrites, preparation, 147, 152. 
 
 Nitrocellulose, preparation, 220; 
 use, 234. 
 
 Nitrogen, occurrence in air, 
 weight over square foot of 
 earth; in all living bodies, 
 source of, 138, 152; diagram 
 of course in nature; prepa- 
 ration; properties, 139; oxides 
 of, 147; essential element in 
 soil, 297. 
 
 Nitrogen dioxide, as catalytic 
 agent, 121; as oxidizing agent, 
 128; from nitric acid; pro- 
 perties, 144; formation "form- 
 ula, polymer, dissociation, 150, 
 153; products on solution, 
 153. 
 
 Nitrogen oxides, to illustrate 
 law of multiple proportions, 
 50. 
 
 Nitrogen tetroxide, reaction with 
 water, 151. 
 
 Nitroglycerine, manufacture by 
 sulfuric acid, 125; manufac- 
 ture, how exploded, 228. 
 
 Nitrous acid, preparation, de- 
 composition, 147. 
 
 Nitrous anhydride, preparation, 
 147, 179; how formed, dissocia- 
 tion, 150, 153. 
 
 Nitrous oxide, preparation, prop- 
 erties, 147; use, 153. 
 
 Noble gases, properties, list of, 
 157; relation to halogens and 
 alkali metals, 158; zero group, 
 list of, 160. 
 
 Nomenclature, binary com- 
 pounds, 15, acids and salts, 
 84, 85. 
 
 Non-coking coals, properties, 192. 
 
 Normal salts, definition, 79, 
 118. 
 
 Normal salts of strong acids are 
 neutral; of weak acids, 
 frequently alkaline, 76. 
 
 Normal sulfates of strong bases, 
 neutral, 125. 
 
 Number of molecules, in cubic 
 centimeter of gas; in equal 
 volumes, 33. 
 
 Oil of vitriol, name for sulfuric 
 
 acid, 285. 
 
 Oleic acid, in fats, 227. 
 Olein, derivation, 227. 
 Opal, form of silicon dioxide, 
 
 containing water, 238. 
 Open hearth steel, characteristics 
 
 of, 333; manufacture, 327. 
 Ores of common metals, reduc- 
 tion of, 277. 
 
 Organic matter, source of car- 
 bonic acid, 62. 
 Orpiment, source of arsenic tri- 
 
 sulfide, 179. 
 Orthophosphoric acid, salts of; 
 
 tribasic, 174, 183. 
 Osmium, as catalyser in forming 
 
 ammonia, 140. 
 -ous, meaning of, 84. 
 Oxalic acid, in preparation of 
 
 carbon monoxide, 211, 216. 
 Oxidation, by action of water and 
 
 air, 14; slow, cause of bodily 
 
 temperature, 13; by bacteria, 
 
 14. 
 Oxidation reduction, opposite 
 
 processes, 24. 
 
360 
 
 INDEX 
 
 Oxides of metals, solution in 
 hydrochloric acid, 82; of anti- 
 mony, 181; of nitrogen, list 
 of, 147; of lead, list; properties, 
 248; of iron, zinc and tin, 
 how reduced, 277. 
 
 Oxidizing agents, action on hy- 
 drochloric acid ; list given ; 82, 
 90; potassium dichromate as, 
 317; in preparation of potas- 
 sium manganate, 321; use of 
 potassium permanganate, 322. 
 
 Oxone, composition, 207. 
 
 Oxy -acetylene flame, use of, 27. 
 
 Oxygen, preparation from red 
 oxide of mercury, 7 ; from potas- 
 sium chlorate, 9, 87; from liq- 
 uid air, 10, 17; properties, 10, 
 17; occurrence, 12; velocity of 
 molecules, 34; weight of liter 
 of, 43; per cent, in air, 57; for- 
 mula of, how determined, 136; 
 in air, how determined, 155, 
 160; preparation from barium 
 peroxide, 268, 275. 
 
 Oxygen acids of chlorine, 84; 
 salts of, 85. 
 
 Oxyhydrogen blowpipe, why tem- 
 perature is limited, 25; use of, 
 26. 
 
 Ozone, preparation, properties, 
 15; allotropic form of oxygen, 
 17; for purification of water, 
 70; weight of liter of, 137. 
 
 Palmatin, derivation, 227. 
 Palmitic acid, in fats, 227. 
 Paper, manufacture, 220. 
 Paraffin, product from petroleum, 
 
 199. 
 Paris green, composition; use, 
 
 179, 1S3. 
 
 Parts, by weight and atoms, 52. 
 
 Peat, properties, 191, 192. 
 
 Pectin, preparation ; use, 224, 234. 
 
 Pectose, presence in fruits, 224, 
 234. 
 
 Per-, meaning of, 49. 
 
 Perchloric acid, properties, 89; 
 use for detection of potassium, 
 90. 
 
 "Perfect" gases, 132. 
 
 Periodic system, discussion of 
 tables, 102; relation to atomic 
 weights, 133; tables, 164, 165. 
 
 Periods of periodic system, 
 162. 
 
 Perkin, W. H., discoverer of 
 mauve, 233. 
 
 Permanent hardness of waters, 
 263. 
 
 Petrolatum, see vaseline, 199. 
 
 Petroleum, composition; occur- 
 rence, 198; use; distillates, 198; 
 products from, 199, 208. 
 
 Phenol, use of benzene in manu- 
 facture of, 202; preparation; 
 use, 228, 235. 
 
 Phenol phthalein, as an indicator, 
 76. 
 
 Phosphates, mineral, use for 
 fertilizers, 125; use as fertilizer, 
 170. 
 
 Phosphine, preparation; proper- 
 ties; comparison with ammonia, 
 172, 183. 
 
 Phosphonium iodide, how pre- 
 pared; comparison with am- 
 monium iodide, 173. 
 
 Phosphoric acid, preparation, 171 ; 
 ionization of, 175; hydrolysis 
 of salts of, 175; by hydrolysis 
 of phosphorus pentachloride, 
 176. 
 
INDEX 
 
 361 
 
 
 Phosphoric acids, list of, 174, 
 183. 
 
 Phosphoric anhydride, see phos- 
 phorus pentoxide, 173. 
 
 Phosphorous acid, by hydrolysis 
 of phosphorus trichloride, 176; 
 dibasic, properties, 176, 183. 
 
 Phosphorus, burning in oxygen, 
 11; kindling temperature of, 
 12; preparation, 171; occur- 
 rence in combination, 170; 
 properties; allotropic forms of, 
 171 ; in manufacture of matches, 
 poisoning by, 171; occurrence; 
 preparation; forms, 182; es- 
 sential element in soil, 297. 
 
 Phosphorus chlorides, hydro- 
 lysis of, 176. 
 
 Phosphorus pentachloride, prepa- 
 ration, 83; preparation, hy- 
 drolysis, 176, 183. 
 
 Phosphorus pentoxide, from com- 
 bustion experiment; properties, 
 11; gives acid with water, 17; 
 preparation; properties; use as 
 drying agent, 173; preparation; 
 use as drying agent, 182. 
 
 Phosphorus trichloride, prepara- 
 tion, 83. 
 
 Phosphorus trichloride, prepara- 
 tion, hydrolysis, 176, 183. 
 
 Photographic plates, effects of 
 uranium and radium upon, 
 268. 
 
 Photography, use of silver halides 
 in, 311. 
 
 Physical science, definition, 1. 
 
 Physics, definition, 1, 5. 
 
 Pig iron, manufacture, 324, 
 333. 
 
 Plants, what growth depends 
 upon, 231. 235. 
 
 Plaster casts, 264. 
 
 Plaster of Paris, how made, 110; 
 manufacture; use, 264; source 
 of, 261; properties, 275. 
 
 Platinic chloride, preparation, 
 332. 
 
 Platinum, melting point, 26; 
 as catalytic agent, 119; action 
 of* aqua regia upon, 146; 
 use; price; as catalyst; aqua 
 regia as solvent of, 331, 332; 
 occurrence; use, 334. 
 
 Plumbic acid, salt of, 248. 
 
 Polymer, definition, 119. 
 
 Porcelain, 257. 
 
 "Positive" in photography ex- 
 planation, 312. 
 
 Positive and negative valences, 
 104, 106. 
 
 Potash salts, where found, 68; 
 occurrence; use, 298. 
 
 Potassium, used to decompose 
 water, 19; detection by per- 
 chloric acid, 90; solubility of 
 salts of, 280; occurrence; es- 
 sential element in soil, 297, 307;. 
 preparation ; properties, 298, 
 306. 
 
 Potassium aluminium sulfate; 
 action in bread making, 230. 
 
 Potassium argenticyanide, use, 
 216, 217. 
 
 Potassium bicarbonate, use, 299, 
 306. 
 
 Potassium bromide, use in medi- 
 cine, 94. 
 
 Potassium carbonate, solution of, 
 alkaline, 213; in preparation of 
 potassium ferrocyanide, 215; 
 from wood ashes by leaching;, 
 use in soap making, 297; 
 properties, 299; preparation, 
 
362 
 
 INDEX 
 
 306; use in preparation of 
 potassium chromate, 317. 
 
 Potassium chlorate, use for pre- 
 paring oxygen, 9, 87; chloric 
 acid from, 88; to prepare potas- 
 sium perchlorate, 89; products 
 when heated, 90; use in flash 
 light powders, 283; prepara- 
 tion, use, 299, 306. 
 
 Potassium chloride, in manufac- 
 ture of gunpowder, 279. 
 
 Potassium chloroplatinate, prepa- 
 ration; use as laboratory test, 
 332. 
 
 Potassium chromate, preparation, 
 properties, 317, 319. 
 
 Potassium cyanide, preparation; 
 use in extracting gold, 215, 217; 
 in preparation of potassium 
 ferrocyanide, 215. 
 
 Potassium dichromate, to oxidize 
 hydrochloric acid, 83; use in 
 manufacture of chrome yellow, 
 250; preparation, properties, 
 use, 317; use in tanning, 318, 
 319. 
 
 Potassium f erricyanide, prepara- 
 tion, use, 216. 
 
 Potassium ferrocyanide; prepa- 
 ration, properties, 215, 217; 
 electrolysis of, 216. 
 
 Potassium fluoride, use in pre- 
 paring fluorine, 96. 
 
 Potassium hydroxide, by de- 
 composition of water, 26; in 
 reversible reaction, 86; prepa- 
 ration, 297, 299; use in prepara- 
 tion of potassium chromate, 
 317. 
 
 Potassium hypochlorite ; products 
 when warmed, 90; how formed, 
 299. 
 
 Potassium iodide, use of, 95, 98. 
 
 Potassium manganate, prepara- 
 tion, 321. 
 
 Potassium nitrate, in soil, source 
 of, 138; from decay of organic 
 matter; from sodium nitrate, 
 144; preparation for manufac- 
 ture of gunpowder, 279; occur- 
 rence; preparation; use; 299, 
 306. 
 
 Potassium oxide, preparation, 
 298. 
 
 Potassium perchlorate, from pot- 
 assium chlorate ; properties, 
 89, 90. 
 
 Potassium permanganate, to 
 oxidize hydrochloric acid, 83; 
 preparation, 321; use, 322. 
 
 Potassium sodium tartrate, prod- 
 uct of action of baking pow- 
 ders, 230. 
 
 Potassium sulfate, use in manu- 
 facturing alum, 257. 
 
 Potassium tartrate, acid, 229. 
 
 Potential, difference between 
 metals and solutions, 242; 
 difference between metals, 243; 
 difference in storage batteries, 
 249. 
 
 Potential energy, 2. 
 
 Pound, grams in one, 57. 
 
 Precipitates, formation of, rever- 
 sible, 113, 127. 
 
 Prefixes, in chemical nomencla- 
 ture, 84, 85, 90; hypo-, 85; per-, 
 49, 85;thio-, meaning, 126, 296; 
 to distinguish salts of ortho- 
 phosphoric acid, 174. 
 
 Pressure of air, at sea level, 36; 
 standard conditions of, 37. 
 
 Principle of van't Hoff-Le Chat- 
 elier, 120. 
 
INDEX 
 
 363 
 
 Priestly, preparation of oxygen 
 gas by, 7. 
 
 Producer gas, manufacture; use; 
 composition, 204, 208. 
 
 Proteins, definition; occurrence, 
 230; effect of digestion upon, 
 231. 
 
 Prussian blue, use in preparation 
 of chrome green, 319. 
 
 Prussic acid, preparation ; poison- 
 ous quality, 215, 217. 
 
 Puddling process, for making 
 wrought iron, 326, 333. 
 
 Pure substances, definition ; 3, 5. 
 
 Purification, by crystallization 
 and distillation, 2. 
 
 Pyrite, 110. 
 
 Pyroboric acid, as source of 
 borax, 253. 
 
 Pyrophosphoric acid, 174; prepa- 
 ration of salts of, 175. 
 
 Qualitative analysis, explana- 
 tion; separation of metals 
 with insoluble chlorides, 335; 
 groups of, 112; separation of 
 metals with sulfides insoluble 
 in dilute acids, 336; detec- 
 tion of sulfates, chlorides, and 
 nitrates, 337. 
 
 Quantitative analysis, explana- 
 tion, method of procedure, 
 337. - 
 
 Quartz, varieties of, 238. 
 
 Quinine, source; use, 233. 
 
 Radioactive elements, disinte- 
 gration of atoms of, 163; 
 detected by gold leaf electro- 
 scope, 269; estimate of life 
 of, 271; three series of; half- 
 life period of, 275. 
 
 Radiolite watch dials, cause of 
 phosphorescence, 269; meso- 
 thorium for, 273. 
 
 Radium, products of decomposi- 
 tion of, 158; discovery; prop- 
 erties, occurrence, 268; price 
 of, 269; products of, 272; table 
 of derivatives, 272 ; how formed ; 
 disintegration of, 275. 
 
 Radium sulfate, properties, 260. 
 
 Raleigh, discovery of argon, 157. 
 
 Ramsay, discovery of argon, 157; 
 of helium and niton, 158. 
 
 Ramsay and Soddy, kelium from 
 radium, 272. 
 
 Reactions, explanation of rever- 
 sible, 24, 80; effect of removing 
 one product upon, 115; endo- 
 thermic, explanation, 201. 
 
 Reduction, definition, 24; of 
 copper oxide, 24, 47; of ores of 
 common metals, 277. 
 
 Red lead, preparation; use; 248. 
 
 Red phosphorus, properties, use, 
 171, 182. 
 
 Remedies, for burns by sulfuric 
 acid, 124. 
 
 Respiration calorimeter, use of , 13. 
 
 Reversible reactions, explana- 
 tion, 24, 80; how expressed, 
 53; effect of removing one- 
 of the products of, 85; in 
 formation of precipitates; in 
 formation of ions, 114; in pre- 
 cipitation of lead sulfide, 114;. 
 effect of temperature upon, 
 120. 
 
 Rochelle salt, compositions; 
 use in preparing cuprous oxide, 
 309. 
 
 Rock crystal, form of silicon 
 dioxide, 238. 
 
364 
 
 INDEX 
 
 Roll brimstone, definition, 110. 
 
 Rose quartz, form of silicon 
 dioxide, 238. 
 
 Rubies, composition, 255. 
 
 Rust, nature; cause of, 284. 
 
 Rutherford, disintegration of ele- 
 ments, 272. 
 
 Safety lamp, 197. 
 
 Safety plugs, composition, 248. 
 
 Sal soda, preparation, 291. 
 
 Saleratus, see Potassium bi- 
 carbonate, 299. 
 
 Salt, how deposited, 68; composi- 
 tion, occurrence, 72; formula, 
 how determined; definition, 75; 
 see sodium chloride, 290. 
 
 Saltpeter, source of nitric acid, 
 143; preparation for manufac- 
 ture of gunpowder, 279; source 
 of, 291; occurrence; prepa- 
 ration, use, 299. 
 
 Salts, definition, 21; acid, how 
 formed; normal, neutral, and 
 acid, defined, 76, 77; normal 
 of strong acids, are neutral; 
 of weak acids, frequently aka- 
 line 76; why called normal; 
 why called acid, 78; neutral; 
 alkaline or acid, 79; from oxy- 
 gen acids of chlorine, 85; o 
 sulfurous acid, preparation; 
 properties, 117; normal, defini- 
 tion, 118; of orthophosphoric 
 acid, 174; of phosphoric acid, 
 hydrolysis of, 175; ammonium, 
 preparation, 191; of metals, 
 how prepared, reversible re- 
 actions involved, explanation, 
 278; rules for solubility of, 280; 
 reactions for preparation of, 
 conditions necessary ; when in- 
 
 soluble, 281; of alkali metals, 
 properties, 290; list of am- 
 monium, 302; of copper, 309; 
 of iron, 329. 
 
 Sand, composition, 239. 
 
 Saponification, defined, 227. 
 
 Sapphire, form of corundum, 255. 
 
 Saturated compounds, definition, 
 199. 
 
 Saturated solutions, 65. 
 
 Scale, in utensils, cause of, 62; 
 in boilers, cause of, 264. 
 
 Science, definition, 1, 5. 
 
 Sedimentary rocks, composition 
 of, formation, 255. 
 
 Selenite, properties, 261. 
 
 Selenium, properties ; com- 
 pounds, use, 126, 128. 
 
 Shales, composition of, forma- 
 tion, 255. 
 
 Sicily, sulfur from, 109. 
 
 Siderite, ore of iron, 324. 
 
 Silica, in manufacture of phos- 
 phorus, 171. 
 
 Silica, occurrence in soil, 237. 
 
 Silicates, occurrence in soil and 
 rocks, 237; occurrence; deriva- 
 tion; list of; kinds in glass, 239; 
 list of, 251; of aluminium, etc., 
 rocks composed of, 254. 
 
 Silicic acid, from soluble glass, 
 240. 
 
 Silicon, occurrence in crust of 
 earth, 237, 238; preparation; 
 properties; use, 238; occur- 
 rence, 251. 
 
 Silicon dioxide, forms of, 251. 
 
 Silver, occurrence ; separation ; 
 use; properties, 310; percen- 
 tage in U. S. coins, 311; sepa- 
 ration, properties, 314. 
 
 Silver arsenate, 179. 
 
INDEX 
 
 365 
 
 Silver arsenite, salt of arsenious 
 acid, 179. 
 
 Silver bromide, use of, 94. 
 
 Silver bromide, properties ; prepa- 
 ration; use in photography, 
 311, 314. 
 
 Silver chloride, preparation ; 
 properties, 311. 
 
 Silver iodide, properties; prepa- 
 ration, 311. 
 
 Silver nitrate, preparation; 
 properties, use, 311, 314. 
 
 Slag of blast furnace, composi- 
 tion, '325. 
 
 Slaked lime, in soda lime, 197; 
 preparation; use, 262; use in 
 Solvay process, 293; use in 
 manufacture of sodium hy- 
 droxide, 294. 
 
 Slow oxidation, bodily tempera- 
 ture maintained by, 13. 
 
 Small calorie, definition, 58. 
 
 Smokeless powder, from gun 
 cotton, 125; basis of, 220. 
 
 Smoky quartz, form of silicon 
 dioxide, 238. 
 
 Soap, manufacture, saponifica- 
 tion, 227; forms emulsion with 
 water, 227; forms emulsion 
 with water, 235; action in 
 "hard" water, 263. 
 
 Soap, soft, 297; soft, hard, manu- 
 facture, 297 ; soft, manufacture, 
 306. 
 
 Soapstone, a silicate, 239. 
 
 Soda lime, composition; use, 197. 
 
 Soddy and Ramsay, helium from 
 radium, 272. 
 
 Sodium, used to decompose 
 water, 19; properties, 73; solu- 
 ble glass a silicate of, 239; 
 preparation by electrolysis, 
 
 278; solubility of salts of, 
 280; occurrence; 290, 305; 
 sodium bicarbonate from; di- 
 sodium phosphate from, 291. 
 
 Sodium acetate, in preparation of 
 methane, 197. 
 
 Sodium aluminate, preparation, 
 256. 
 
 Sodium ammonium phosphate, 
 175. 
 
 Sodium antimonite, preparation, 
 181. 
 
 Sodium bicarbonate, for wounds 
 by sulfuric acid, 124; use with 
 calomel, 287; product in diges- 
 tion; function in body, 291; 
 preparation by Solvay process, 
 293. 
 
 Sodium bromide, occurrence ; elec- 
 trolysis, 93; use in medicine, 94. 
 
 Sodium carbonate, solution of, 
 alkaline, 213; to remove per- 
 manent hardness in water, 
 264; source of, 291; prepara- 
 tion, 291; product of Le Blanc 
 soda process; alkalinity of, 
 292; use in manufacture of 
 sodium hydroxide, 294; why 
 hydrolyzed, 305; preparation, 
 305. 
 
 Sodium chloride, electrolysis of 
 fused, of aqueous solution, 72; 
 occurrence; office in digestion; 
 importance in chemical indus- 
 tries, 290, 305. 
 
 Sodium cyanide, preparation; use 
 in extracting gold, 215, 217. 
 
 Sodium hydroxide, by decompo- 
 sition of water; affects red 
 litmus, 20, 26; by electrolysis,. 
 73; alkaline reaction, 73; 
 formed by ionization of phos- 
 
366 
 
 INDEX 
 
 phoric acid, 175; in soda lime, 
 197; solvent for stannic hy- 
 droxide, 246; source of, 291; 
 preparation; use, 294, 305; 
 properties, 295. 
 
 Sodium hydroxide, preparation, 
 305. 
 
 Sodium "hyposulfite," use of 
 term, 126; see sodium thio- 
 sulfate, 296. 
 
 Sodium iodide, occurrence, elec- 
 trolysis, 93. 
 
 Sodium metaphosphate, prepara- 
 tion, 175. 
 
 Sodium nitrate, use in manufac- 
 ture of sulfuric acid, 123; 
 occurrence, use, 144; in manu- 
 facture of gunpowder, 279; 
 occurrence, use, 291, 305. 
 
 Sodium oxide, preparation, 295. 
 
 Sodium peroxide, preparation; 
 use, 296, 305. 
 
 Sodium silicate, preparation ; use, 
 296. 
 
 Sodium stannate, preparation ; 
 use, 246. 
 
 Sodium sulfate, in Le Blanc 
 soda process, 292. 
 
 Sodium sulfide, in Le Blanc 
 soda process, 292. 
 
 Sodium sulfite, preparation, 117; 
 preparation; use, 296, 305. 
 
 Sodium sulfite, acid, preparation, 
 116; use, 117. 
 
 Sodium thiosulfate, how formed, 
 properties, use, 125, 128; prepa- 
 ration, why so named, use, 
 296; use to dissolve salts of 
 silver, 311, 312, 
 
 Soft soap, 297. 
 
 Soils, essential elements in, 
 297. 
 
 Solder, composition, 245; alloy 
 of lead, 248. 
 
 Solubility, in hot and cold water, 
 65; of salts, general rules for, 
 280. 
 
 Soluble glass, preparation; prop- 
 erties, use, 240, 251; prepara- 
 tion; use, 296, 305. 
 
 Solution pressure of metals, ex- 
 planation; cause of electric 
 current between metals, 242, 
 251. 
 
 Solutions, change of freezing and 
 boiling points, 64; chemical 
 activity in; ionization in, 65; 
 supersaturated; saturated, 65, 
 69; neutral defined, 75, 79. 
 
 Solvay process for preparation 
 of sodium bicarbonate, 293. 
 
 Specific gravity of elements, rela- 
 tion to position in periodic 
 system, 167, 168. 
 
 Spectra, how obtained, 302. 
 
 Spectroscope, explanation, 302. 
 
 Spectrum analysis, 302. 
 
 Spectrum, solar, why lines are 
 dark, 303. 
 
 Sphalerite, 110; composition, 284; 
 zinc sulfide, 285; mineral source 
 of zinc, 288. 
 
 Spiegeleisen, use in manufac- 
 ture of steel; preparation, 
 321. 
 
 Spontaneous combustion, how 
 caused; how avoided, 14. 
 
 Stannic acid, in fireproofing cot- 
 ton goods, 246; see stannic 
 hydroxide, 246. 
 
 Stannic chloride, preparation; 
 properties; use, 245. 
 
 Stannic hydroxide, preparation; 
 properties, use, 246, 251. 
 
INDEX 
 
 367 
 
 Stannic oxide, in fireproofing 
 cotton goods, 246. 
 
 Stannous chloride, preparation; 
 properties; use as reducing 
 agent, 245. 
 
 Starch, occurrence; description; 
 use; granules of, 221; test for 
 iodine, 222; glocuse from, 223; 
 occurrence, 234. 
 
 Stearic acid, in fats, 227. 
 
 Stearin, derivation, 227. 
 
 Steel, characteristics of; Bes- 
 semer; open hearth, 326; what 
 hardness depends upon, 333. 
 
 Steel furnace, open hearth, tem- 
 perature of, 26. 
 
 Stereotype metal, use of bismuth 
 in, 181. 
 
 Stibine, preparation; properties, 
 180, 184. 
 
 Stibnite, source of antimony; 
 properties, 179. 
 
 Still, to concentrate alcohol, 225. 
 
 Storage batteries, construction, 
 248; action of; charging of; 
 difference in potential between 
 plates, 248; exhaustion, how 
 tested, 249; substances formed 
 in, 252. 
 
 Storage of gases, 9. 
 
 Stove polish, graphite, 187. 
 
 Strontium, occurrence; proper- 
 ties, 268. 
 
 Strychnine, properties, 233. 
 
 Sublimation, definition, 94. 
 
 "Subnitrate" of bismuth, see 
 bismuth, basic nitrate, 182. 
 
 Substances, pure, definition, 2; 
 not volatile, change in freezing 
 and boiling points by, 64. 
 
 Sucrose or cane sugar, prepara- 
 tion, 222, 234. 
 
 Suffixes, in chemical nomencla- 
 ture, 90; -ate, 85; -ide, 15; 
 -ic, 84; -ite, 85; -ous, 84. 
 
 Sugar, kinds of; preparation, 222. 
 
 Sugar of lead, see lead acetate, 
 250. 
 
 Sulfantimonite, ammonium, 181. 
 
 Sulfarsenite, ammonium, 179. 
 
 Sulfates, list of insoluble, 125, 
 128. 
 
 Sulfides in analysis, 112; pre- 
 cipitation due to ions, 113; 
 effect of strength of acids in 
 precipitation of, 115; of zinc 
 and antimony, how reduced, 
 277. 
 
 Sulfides of copper and lead, how 
 reduced, 278. 
 
 Sulfur, burning: in oxygen, 11; 
 occurrence,, how secured, 109; 
 occurrence as sulfates and 
 sulfides r 110'; use in sulfuring 
 fruit; in lime-sulfur wash; in 
 gunpowder; in mamuf acture ; 
 forms of r 110;, use in dyes, 111; 
 allotropie- forms,. 111 ;liquid and 
 gaseous r properties y 111; boil- 
 ing point,. 112; occurrence; use; 
 allotropie forrns r 127. 
 
 Sulfur dioxide^ by burning sul- 
 fur in oxygen; properties, 11; 
 gives acid with water, 17; 
 use in sulfuring fruit, '110; 
 occurrence ; properties ; prepara- 
 tion, 115; laboratory prepara- 
 tion of r 118; preparation from 
 iron pyrites, 122; preparation; 
 use r 127; preparation, 145. 
 
 Sulfur trioxide, formation from 
 sulfur and oxygen reversible; 
 preparation; properties, 118, 
 119 r 
 
368 
 
 INDEX 
 
 Sulfuric acid, with zinc; with iron, 
 21; electrolysis of, 42; use to 
 liberate chlorine, 87 ; from iron 
 pyrites, 110; manufacture by 
 contact process, 119; prepara- 
 tion by lead chamber process, 
 121. 
 
 Sulfuric acid, concentration of; 
 properties, 123; charring by; 
 wounds by, remedies; use of, 
 124; use in preparation of 
 nitric acid, 144; as a catalyzer, 
 211; behavior in storage bat- 
 teries, 249; use in Le Blanc 
 soda process, 292. 
 
 Sulfuric anhydride, see sulfur 
 trioxide, 120. 
 
 Sulfuring fruit, with sulfur diox- 
 ide, 115. 
 
 Sulfurous acid, preparation, 
 1 17 ; dibasic ; salts of, 117; prepa- 
 ration; properties, 128. 
 
 "Super-phosphate" of calcium, 
 composition, 265 ; use as fer- 
 tilizer, 275. 
 
 Supersaturated solutions, 65. 
 
 Symbols of elements, 51, 56. 
 
 Sympathetic ink, use of cobalt 
 chloride in, 331, 334. 
 
 Synthesis analysis, explanation, 
 42, 55. 
 
 Synthesis, definition, 3; of com- 
 pounds, 5; of ammonia, 142; 
 of nitric oxide, 148. 
 
 Talc, a silicate, 239. 
 
 Tan bark, use in manufacture of 
 white lead, 250. 
 
 Tannin, source; use in tanning, 
 318. 
 
 Tanning, chrome, use of potas- 
 sium dichromate ;n, 317, 319. 
 
 Tar distillates, source; use, 189. 
 
 Tartaric acid, preparation ; source ; 
 salts, 228, 235. 
 
 Tellurium, properties, occurrence, 
 126, 128; exception in periodic 
 table, 168. 
 
 Temperature, lowest obtained, 
 35; how measured, 34, 40; 
 absolute, explanation, 34, 40; 
 standard conditions of, 37; 
 influence on rate of combina- 
 tion of elements in reversible 
 reaction, 120. 
 
 Temporary hardness of waters, 
 263. 
 
 Tetraphosphorus trisulfide, use 
 in the manufacture of matches, 
 172, 182. 
 
 Thermite.process, Goldschmidt's, 
 256, 258. 
 
 Thermometers, graduation of, 
 34; use of mercury in, 286. 
 
 Thio-, meaning, 126, 296. 
 
 Thorium, series of derivatives, 273. 
 
 Thyroid gland, iodine in, 95. 
 
 Tin, occurrence; preparation; 
 properties; use, 241, 251 ; alloys 
 of, description, 245; alloys of, 
 list of; chlorides of, 251; re- 
 duction of oxides of, 277. 
 
 Tinware, cause of rust, 243. 
 
 Toluene, derivative of coal tar r 
 189; use in making trinitrotol- 
 uene (T. N. T.) 202. 
 
 Toxins, how formed; properties, 
 232, 235. 
 
 Tri-, prefix, 174. 
 
 Tribasic acid, defined, 76, 118. 
 
 Tricalcium phosphate ; occur- 
 rence; 170. 
 
 Trinitrotoluene (T. N. T.), prepa- 
 ration; u^e,, 202. 
 
INDEX 
 
 369 
 
 Trisilver arsenate, 179. 
 Tungsten, why suggested for 
 
 electric lights, 168; discovery 
 
 of use, 318, 319. 
 
 Type-metal, composition of, 180. 
 Typhoid fever, from impure water, 
 
 64. 
 
 Ultra-violet light for purification 
 of water, 70. 
 
 Unit volume, measure of, 132. 
 
 Unit volume of gases, how deter- 
 mined, 132. 
 
 Unsaturated compounds, defini- 
 tion, 199. 
 
 Uranium, as catalyser in forming 
 ammonia, 140; effect of com- 
 pounds on photographic plates, 
 268. 
 
 Valence, negative in halogens, 99; 
 explanation, 101; nomencla- 
 ture of, 102; varying, 103; 
 positive and negative, 104, 
 107; table of, 106; table of 
 common, 107; definition; ex- 
 pressed by graphical formulas, 
 107; indicated by position of 
 element in periodic system, 
 163. 
 
 Van't Hoff-Le Chatelier, principle 
 of, 120, 128; application of 
 principle of, 120, 142, 148. 
 
 Vapor pressure of water, 60; 
 explanation, 62; law of; table 
 for ice and water, 63. 
 
 Vaporization of water, latent heat 
 of, 59, 69. 
 
 Vaseline, product from petroleum, 
 199. 
 
 Velocity of molecules, law of, -33; 
 of hydrogen; of oxygen, 33, 
 
 24 
 
 Ventilation, standard of, 159. 
 
 Vinegar, product of acetic fer- 
 mentation, 227, 235; method 
 of manufacture, 227. 
 
 Vitriols,compounds so named,285. 
 
 Volatile compounds, conditions 
 for formation of, 278. 
 
 Volume of gas, variation with 
 temperature, 35; reduction of 
 gases to standard, 37; gram 
 molecular weight of; measure 
 of, 134. 
 
 Volumes of gases, number of 
 molecular in same, 33, 40; 
 law of change of; method of 
 finding, 37, 40; ratio in simple 
 numbers, 132. 
 
 Washing soda, see sodium car- 
 bonate, 291. 
 
 Water, decomposition of; by red 
 hot iron, 19, 26; by sodium or 
 potassium, 19; by electricity, 
 44. 
 
 Water, composition by weight, 
 44, 45, 50; volumetric compo- 
 sition of, 44; supercooled, 58; 
 boiling of; superheated; latent 
 heat of vaporization of, 59; 
 expansion on freezing, 59; 
 cooling by convection; vapor 
 pressure of, 60, 69; natural, 61; 
 "hard," 62; table of vapor pres- 
 sures, 63; as disease carrier, 
 64, 69; mineral, 67; latent heat 
 of vaporization of; maximum 
 density, 69; ionization, 74; 
 methods of purification, 70; 
 hard, explanation, 263; tem- 
 porary and permanent hard- 
 ness, explanation; how re- 
 moved, 264. 
 

 Water gas, preparation; proper- 
 ties; poisonous quality, 203 r 
 208. 
 
 Water vapor in air, 156. 
 
 Weight, no change of in burning, 
 7. 
 
 White lead, preparation; prop- 
 erties; use, 250.; composition; 
 use, 252. 
 
 White vitriol, name for zinc 
 sulfate, 285. 
 
 Window glass, composition, 239. 
 
 Wire gauze, use in safety lamp, 
 197. 
 
 Wood material, changes during 
 geological time, 192. 
 
 Wood's metal, composition, use, 
 181. 
 
 Wrought iron, manufacture; use, 
 326, 333. 
 
 Xenon, where found, 158. 
 X-rays, effects similar to those of 
 
 radium, 269. 
 X-ray spectra, to determine 
 
 atomic number,, 168. 
 
 Yeast, cause of fermentation, 224. 
 
 Zero,, absolute,. explanation,. 
 34. 
 
 Zero grotrp r noble gases, list of, 
 160. 
 
 Zinc, for preparation of hydrogen, 
 21; as eoating for iron, 243; 
 reduction of oxides of; reduc- 
 tion of sulfides of, 277; occur- 
 rence; preparation; properties; 
 use; alloys of, 284, 288; use in 
 batteries, 285; use in refining 
 silver, 310; use in cyan-'de 
 process for gold, 312; use in 
 nickel silver, 331. 
 
 Zinc chloride, how formed, 21; 
 how prepared, 83; use as 
 disinfectant, 288. 
 
 Zinc hydroxide, properties, 260. 
 
 Zinc oxide, preparation; prop- 
 erties; use; when preferable 
 to white lead in paint, 285, 
 288. 
 
 Zinc sulfate, how formed, 21; 
 popular name of, 288. 
 
 Zinc sulfide, action of radium 
 upon, 269; use of phosphores- 
 cent screen of, 274; occurrence; 
 preparation, 285; color, 285. 
 
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