ESSENTIALS OF CHEMISTRY BY JOHN C. HESSLER, PH.D. t INSTRUCTOR IN CHEMISTRY, THE UNIVERSITY OF CHICAGO; LATE INSTRUCTOR IN CHEMISTRY, THE HYDE PARK (CHICAGO) HIGH SCHOOL AND ALBERT L. SMITH, PH.D. INSTRUCTOR IN CHEMISTRY, THE ENGLEWOOD HIGH SCHOOL, CHICAGO ov TroXX' aXXa BOSTON, U.S.A. BENJ. H. SANBORN & CO. COPYRIGHT, 1902, By JOHN C. HESSLER AND ALBERT L. SMITH SHje JFort SAMUEL USHER 176 TO 184 HIGH STREET BOSTON, MASS. PREFACE. The interest aroused by the introduction of laboratory science into secondary schools, a decade and a half ago, brought out a number of text-books of Chemistry. Al- though most of these books had directions for laboratory work scattered through the descriptive matter, yet the laboratory exercises Avere illustrations of the text proper rather than working notes for the student. A further source of confusion lay in the fact that the experiments capable of being performed by the student were hope- lessly mixed with those intended for the teacher. When the separate " laboratory manual " appeared it was usually characterized by the same faults. Within recent years many new text-books of Chem- istry have been written for secondary schools, but, with a few exceptions, the new books, like the old ones, are impractical. They are either too diffuse in description, the laboratory work being left chiefly to the invention of the teacher, or they are merely laboratory manuals, without enough descriptive matter to make them useful as text-books. Moreover, the laboratory exercises of many modern books are wholly beyond the capabilities of the average student. There is, therefore, a demand for a text-book of Chemistry containing an adequate and scientific account 221752 iv PREFACE. of such of the fundamental facts, laws, and theories of the subject as are adapted to the needs of secondary schools, and also specific directions for the laboratory work, directions that have been tested and found prac- ticable. This demand it is the design of the present book to Lieet. The authors have had exceptional opportunities to know the capacity of the average student and the train- ing of the average teacher of science in secondary schools. They are fully aware of the limitations of these schools both as to laboratory equipment and as to the time which may reasonably be expected for the study of Chemistry. They have prepared the book with these limitations constantly in mind. A connected treatment of the descriptive matter of this work is attained by the division of the book into three parts: (1) the text proper; (2) the laboratory exercises ; (3) the handbook. The text and the labora- tory exercises are bound together ; the handbook is in pamphlet form. The text proper may be characterized by saying that it recognizes the fact that the terms and the ideas of Chemistry are outside of the common experience, and that it is useless to expect the pupil to grasp theoretical con- ceptions before he has become acquainted with the fun- damental phenomena of the science. The arrangement of topics is such that the early chapters of the book are mainly descriptive ; theoretical ideas are not introduced until later, and then only in an elementary manner. PREFACE. V An illustration of the arrangement is the case of molecular masses, atoms, and atomic masses, which are not mentioned at all until Chapter XVI, after the ele- ments hydrogen, oxygen, chlorine, nitrogen, sulphur, and carbon, and their most important compounds, have been studied. Equations, although introduced early to accustom the pupil to the fact that chemical reactions are quantitative, are for some time (up to page 69) written out in full rather than in symbolic form. The idea of equilibrium is introduced first in connec- tion with the study of diffusion ; later its application is extended to chemical reactions. The aim has been to make the text modern and scientific, yet not too difficult for secondary students. The laboratory exercises are placed together after the text. The directions for these exercises are specific. The quantities of reagents to be employed have been carefully considered. TJie apparatus required is simple and within the reach of every school. The experiments are so arranged that they may be used in schools in which only one-hour laboratory periods are possible, as well as in those able to give two or more hours of con- secutive time. The experiments are mainly qualitative, as experience has shown that they must be, even with beginning college students. Several experiments of a quantitative nature have, however, been introduced; for these only the simplest apparatus is required. Vi PREFACE. In the earlier experiments, especially, the directions are very explicit; the teacher should, therefore, be able to use his laboratory time from the beginning in question- ing individual students and in directing their work at close range. Much material designed chiefly for the teacher's use has been put into the handbook. The handbook con- tains, also, a list of experiments to be performed by the teacher before the class. These experiments are care- fully planned, and the directions for carrying them out, accurate. Every teacher will thus have at his command a series of demonstrations with which to supplement the laboratory exercises performed by the students. The authors suggest that the several parts of this book be used as follows : (1) The student is to perform the experiments for the week in the laboratory, taking notes upon his work. These notes are to be examined weekly by the te: cher. (2) The teacher is to perform the demonstration ex- periments for the week, the class taking notes. (3) The teacher is to assign topics for recitation. These topics should demand (a) an account of the laboratory work, (>) a description of the teacher's ex- periments, ( . . 49 XXXVII. Carbon 50 XXXVIII. Carbon Dioxide, I 51 XXXIX. Carbon Dioxide, II 52 XL. Reduction. Effect of Heat on Carbonates ... 53 XLI. Flames 54 XLII. Weight of a Liter of Oxygen 55 XLIII. Bromine 57 XLIV. Iodine and Hydriodic Acid 58 XLV. Comparison of the Halogen Acids 60 XL VI. Hydrogen Peroxide 61 XLVII. Phosphorus and Phosphoric Acid 62 XLVIII. Arsenic 63 XLIX. Antimony 65 L. Bismuth 66 Li. Borax and Boric Acid 66, LII. lonization 68 LIII. Hydrolysis and Mass Action 68 LIV. Sodium Compounds 69 LV. Potassium Compounds 71 LVI. Solubility of Potassium Chloride 72 LVII. Ammonium Amalgam. Distinctions between the Alkali Metals 73 LVIII. Calcium 74 LIX. Water of Crystallization in Gypsum 75 LX, Strontium and Barium .... ,76 CONTENTS. XXI EXPERIMENT. PAGE. LXI. Water of Crystallization in Barium Chloride . . 77 LXII. Magnesium 78 LXIII. Zinc 79 LXIV. Equivalent of Zinc 81 LXV. Cadmium 81 LXVI. Mercury 82 LXVII. Copper 84 LXVIII. Silver 85 LX1X. Aluminum 86 LXX. Iron 87 LXXI. Nickel and Cobalt 90 LXXII. Manganese Compounds 90 LXXIII. Chromium Compounds 91 LXXIV. Lead 93 LXJLV. Tin ....... 94 CHEMISTRY. INTRODUCTION. 1. Definition of Chemistry. Chemistry may be defined as the science which deals with the different kinds of matter and their transformations. The kinds of matter are called substances. Many substances occur ready- formed in the earth ; but many others are not so found, and must be made from existing bodies. Chemistry, therefore, consists largely of a description of the ways in which substances occur in nature, of the methods by which they may be produced in the laboratory, and of the properties, or characteristics, by which they may be distinguished from one another. Descriptive Chemistry alone, however, cannot give a connected and an intelligent view of the whole science ; this can result only from a study both of the laws gov- erning chemical phenomena, and of the most important theories that men have suggested to explain the laws. 2. Importance of Chemistry. --The ideas of modern Chemistry are of great importance, both for their own 2 INTRODUCTION. sake, and because they add so much to other depart- ments of knowledge ; no apology is therefore needed for the presence of Chemistry in a course of study. We may, however, summarize in brief form the reasons why every student in a secondary school should master at least the elements of chemical science : 1. A good course of laboratory work in Chemistry disciplines the mind, as few courses can, in independent and honest observation of phenomena. 2. A knowledge of this science is necessary for the intelli- gent study of other natural sciences, such as geology, astronomy, biology, physiology, etc., which make special application of its general ideas. 3. Chemistry is intensely practical, for its facts are in the most common use in the arts and in e very-day life. As illustrations of the practical character of chemical knowl- edge we may cite its application in medicine, in sanitation, in domestic science, in the extraction of metals from their ores, in the refining of petroleum, and in the manufacture of steel, illuminating and fuel gas, paints, dyestuffs, food products, ice, alcohol, soap, glass, paper, explosives, etc. To be sure, it is impossible for an elementary course in Chemistry to give all the facts relating to such topics as those suggested above ; but it does give the facts and the theories that are fundamental, and by means of which even more com- plex phenomena must be interpreted. 3. Relation between Chemistry and Physics. - Chemistry and Physics are closely related sciences, for both together have as their object the study and the \ "' ' .Vli-'IU^ I ; & ':- INTRODUCTION. 3 explanation of the general phenomena, or changes, that take place in the material universe. They are, in fact, two points of view from which natural phenomena may be considered. Physics has to do chiefly with transformations of energy, and with matter only as that upon which energy acts to produce phenomena. Uhemistry, on the other hand, is largely concerned with the various forms of matter, and with energy-changes only as they result in the formation of new sub- stances. Accordingly, phenomena are usually distinguished as either physical or chemical phenomena. To the former class belong those changes in which the substance, in connection with which the energy manifests itself, is not permanently altered, but regains its original properties. A physical change may, therefore, be repeated with the same substance after the substance has resumed its former condition. Illustrations are : the magnetization of a knife-blade; the production of light by means of white-hot iron ; the melting of ice, and the vaporization of water. Chemical phenomena, on the other hand, involve a permanent alteration of the properties of the substances used. Thus, a piece of burnt magnesium does not assume its original condition on cooling ; rusted iron is no longer iron ; carbonic acid gas is neither carbon nor oxygen, although these two substances were put together to produce it. In the same way we distinguish between the properties INTRODUCTION. of substances, calling those properties physical which require only physical phenomena for their exhibition, and those properties chemical which are capable of being manifested only by some chemical change. Thus the color, specific gravity, melting point, crystalline form, solubility, etc., of sulphur would be termed physical properties of sulphur ; its power of burning in air is, on the contrary, a chemical property. The description of a substance in Chemistry includes its most important physical, as well as its chemical, properties. 4. Reagents and Reactions. Substances which have a chemical effect upon one another are said to react, and a chemical change is therefore called a re- action. The substances which react are called reagents, or factors. Thus, when copper is treated with concen- trated nitric acid, a chemical reaction takes place, and the copper and the nitric acid are the reagents (or factors). The adjective " reagent " is often applied to sub- stances in the form in which they are commonly used in the laboratory. Thus " reagent " ammonia means an aqueous solution of ammonia ; ammonia itself is a gas. Similarly, by " reagent " sodium hydroxide we mean the solution of solid sodium hydroxide in water, this being INTRODUCTION. 5 the form in which sodium hydroxide is most frequently used in the laboratory. 5. Elements and Compounds. Almost all 01 tne substances found in the earth may be shown, by one method or another, to consist of two or more different kinds of matter, and are therefore called compound substances, or compounds. There are, however, between seventy and eighty substances which it is impossible for us to decompose, with our present methods, into simpler kinds of matter ; these substances are, therefore, called elementary substances, or simply elements. Less than half of the substances believed to be elementary are really found free in the earth and its atmosphere ; the others occur only in a combined form. Although the number of elements is so small, the number of compounds they are actually known to form is very large probably not less than one hundred thousand. The number of compounds theoretically pos- sible, but not yet known to exist, is also very large. A list of the substances usually considered elementary follows. The letter, or combination of letters, given after the name of each element is called the symbol of the element. It is not in- tended that all these symbols shall be learned now, but rather by association with the names of the elements as the latter are studied. In cases in which symbols are formed from the Latin (or Greek) names of the elements instead of the English names, both names are given. 6 INTRODUCTION. ELEMENTS. SYMBOLS. ELEMENTS. SYMBOLS. Aluminum. Al. Neodymium. Nd. Antimony. Sb. Nickel. Ni. Argon. A. Niobium. Nb. Arsenic. As. Nitrogen. N. Barium. Ba. Osmium. Os. Beryllium (Glucinum). Be. (Gl.) Oxygen. O. Bismuth. Bi. Palladium. Pd. Boron. B. Phosphorus. P. Bromine. Br. Platinum. Pt. Cadmium. Cd. Potassium (Kallum). K. Caesium. Cs. Praseodymium. Pr. Calcium. Ca. Rhodium. Rh. Carbon. C. Rubidium. Rb. Cerium. Ce. Ruthenium. Ru. Chlorine. Cl. Samarium. Sm. Chromium. Cr. Scandium. Sc. Cobalt. Co. Selenium. Se. Copper (Cuprum). Cu. Silicon. Si. Erbium. Er. Silver (Argeutum). Ag. Fluorine. Fl. Sodium (Natrium). Na. Gallium. Ga. Strontium. Sr. Germanium. Ge. Sulphur. S. Gold (Aurum). Au. Tantalum. Ta. Helium. He. Tellurium. Te. Hydrogen. H. Thallium. Tl. Indium. In. Thorium. Th. Iodine. I. Tin (Stannum). Sn. Iridium. Ir. Titanium. Ti. Iron (Ferrum). Fe. Tungsten (Wolframium). W. Lanthanum. La. Uranium. Ur. ' Lead (Plumbum). Pb. Vanadium. Vd. Lithium. LI. Ytterbium. Yb. Magnesium. Mg. Yttrium. Y. Manganese. Mn. Zinc. Zn. Mercury (Hydrargyrum). Hg. Zirconium. Zr. Molybdenum. Mo. Besides the elements given in the above list, there are several other substances which are considered by some INTRODUCTION. 7 chemists to be elementary, but, the true nature of these substances being still in dispute, their names are omitted. 6. Relative Abundance and Importance of the Elements. The elements are by no means equally abundant ; nor are they all of equal importance for the organic forms existing upon the earth. It is probable that only eleven are absolutely essential to animal life. These are, Carbon, Sulphur, Phosphorus, Oxygen, Calcium, Potassium, Nitrogen, Sodium, Iron. Hydrogen, Chlorine, If four more were present, viz. : Silicon, Magnesium, Aluminum, Fluorine, savage life, upon an earth similar to ours, would be possible. With seven additional elements, viz. : Gold, Tin, Manganese, Silver, Zinc, Mercury, Platinum, modern civilization might exist. The great inequality in the distribution of the ele- ments is shown, in a rough way, by Fig. 1. As there 8 INTRODUCTION. indicated, silicon and oxygen together make up about three fourths of the earth's solid crust. ttf* Potassium l./ . All others not \./<> FIG. i. We shall begin our study of Chemistry in Chapter I with the element that is in some respects the most remarkable of all, the element hydrogen. 7. Exercises. Classify the following as either physical or chemical phe- nomena : The souring of milk, the burning of wood, the evaporating of water, the tarnishing of silver, the dissolving of sugar in water, the bleaching of muslin, the melting of lead. CHAPTER I. HYDROGEN. 8. Existence. Hydrogen is a light, colorless gas that is found in a free condition only in compara- tively small amounts on the earth chiefly in the air and in volcanic gases. It exists in great quantities, however, in the atmosphere of the sun. Although rare in the uncombined form, hydrogen is a constituent of many important and abundant compounds, such as organic substances, water, acids, etc. The most com- mon compound of hydrogen is water. One ninth, by weight, of water is hydrogen, and the remainder is another colorless gas, viz., oxygen. The name "hydrogen" means " a producer of water"; while " oxygen " means " a producer of acids." 9. Common Method of Preparation. Since water is a compound of hydrogen, the decomposition of water will, of course, give hydrogen ; but a much more con- venient way to prepare the gas is to decompose certain acids. Acids, like water, contain hydrogen, and give it up readily when treated with certain metals under ap- propriate conditions. The metal commonly used is zinc, and the acid either hydrochloric or dilute sul- phuric acid, 10 HYDROGEN. Generation and Collection of Hydrogen. The gas is usually produced in a generating flask (Fig. 2) provided with a " thistle " or " safety " tube and a delivery tube reaching into the water of a water pan (called also a " pneumatic trough "). The flask contains zinc. Acid is added through the thistle tube until the lower end of the thistle tube is immersed ; the evolved gas escapes through the FIG. 2. delivery tube and bubbles up through the water in the pan. Here the hydrogen may be collected in appropriate " receiv- ers " (test tubes, bottles, etc.) filled with water and inverted over the end of the delivery tube. Or, since hydrogen is much lighter than air, it may be collected by displacing the air of the receiver, as is shown in Fig. 3. N. B. Apparatus in which hydrogen is being gener- ated must not be brought near a flame ! If the action between metal and acid is not brisk, it PURIFICATION OF HYDROGEN. 11 may be hastened by adding a few drops of copper sulphate solution. The copper sulphate reacts with a portion of the zinc, precipitating copper, which forms a black deposit upon the zinc; thus coated, the latter acts readily upon the acid. The action of zinc and acid re- sults in a very considerable evolution of heat. Self-Regulating Generator. Instead of the ordinary generating flask, a Kipp's or other self-regulating apparatus may be used to supply hydrogen. Kipp's apparatus (Fig. 4) consists of three globes. The upper globe is in communication with the lower globe, and the middle globe with the lower globe, but the upper globe and the middle globe are not connected. The upper and lower globes contain dilute acid, but the middle globe contains zinc. This is the condition of the apparatus when at rest, with the stop- cock closed. When the stopcock is opened, the liquid of the upper globe falls into the lower globe, and the liquid in the lower globe rises into the middle globe, thus displacing the gas of the middle globe, and forcing it out through the stopcock. The acid which enters the middle globe reacts with the zinc, forming more hydrogen, which either escapes through the stopcock, or, if the latter is closed, forces the acid back into the lower globe and thence into the upper globe. The gas in the middle globe is thus ready for instant use. FIG. 3. 10. Purification of Hydrogen. If the metal or the acid used in preparing hydrogen is of " commercial " 12 HYDROGEN, grade, the hydrogen will be impure, as may be inferred from its disagreeable odor. We may remove most of the impurities as well as volatile acid, if hydrochloric acid is used by passing the gas through a mixture of sodium hy- droxide and potassium permanga- nate solutions. To dry it we use some dehydrating agent, e. g., gran- ular calcium chloride. The apparatus for preparing comparatively pure hydrogen is shown in Fig. 5. The hydrogen is generated in a flask (or Kipp's apparatus), and is passed through a bottle containing potassium permanganate dissolved in 10% sodium hydroxide solution, and then through a U-tube of calcium chloride. A sec- ond bottle of the permanganate solu- tion will make the purification more complete. The exit tube of the U- tube is drawn out to a small opening, so that the hydrogen shall issue in a steady stream. If the washing of the hydrogen by the " alkaline perman- ganate " solution has been success- ful, the gas will now be practically odorless. FIG. 4. The stream of gas may be lighted if the following precautions are observed : PURIFICATION OF HYDtiOGEN. 13 In every case, before we light a jet of hydrogen (the same precautions apply to other inflammable gases), we collect a test tube full by displacing the air, and then carry the test tube in this case, mouth down to a gas jet or other flame at least four feet away. The gas in the test tube is thus set on fire, with explosion, if there is still much air mixed with the hy- drogen, but quietly, if the air originally in the apparatus has been displaced. We then carry the test tube back to FIG. 5. the stream of hy- drogen and repeat the operation with freshly collected test tubes full of the gas, until the hydrogen stream is lighted. The reason for all this precaution is that it is unsafe to light a confined mixture of hydrogen and air; and we can be sure that the displacement of the air in the apparatus is complete only when the time needed for the burning of the test tube of hydrogen is greater than that required for the return of the test tube to the jet of hydrogen. The action of zinc upon dilute sulphuric acid gives, 14 HYDROGEN. besides hydrogen, a substance called zinc sulphate. This remains in the solution, but may be obtained from it as a white, crystalline solid. ii. Chemical Properties. The hydrogen flame is almost colorless, but very hot, as holding a piece of platinum in it will show. Indeed, there is enough heat liberated by the burning of a gram of hydrogen in oxy- gen gas to raise the temperature of about 34,000 grams of water 1 C., or about 340 grams from the freezing point to the boiling point. An apparatus for making use of this great heat evolution is the oxyhydrogen blowpipe (Fig. 6). This consists of a small inner tube communicating with a *" *r? ,- ^ tank of oxygen, and a larger outer tube in connection with a tank of hydrogen. Both gases are greatly H compressed. The hydrogen is FIG. 6. first turned on and lighted ; then the oxygen is allowed to escape very slowly. Thus a flame is produced which is so hot that it will melt platinum. (Platinum melts at about 1700 C.) A piece of quicklime held in the oxyhydrogen flame becomes white hot and gives off much light ; this is the so-called calcium, or Zime, light. In the production of the lime light for stereopti- cons, illuminating gas is generally used instead of hydrogen. The ordinary blast-lamp of laboratories is similar to the oxy- hydrogen blowpipe, but the gases used are illuminating gas and ordinary air ; as a result, the temperature produced is by no means as high as that of the oxyhydrogen flame. The hydrogen flame has only one zone, or region, of UNION OF HYDROGEN WITH OXYGEN. 15 combustion. A vertical section of it would have the appearance shown in Fig. 7, - a is the central space of unhurried hydrogen ; b is the region of combustion. The bearing of this fact will be understood when Ave come to study more complex flames. 12. Union of Hydrogen with Oxygen. - The product formed when hydrogen burns in air FIG> 7 is water, as we may prove readily by holding over the burning hydrogen a beaker of cold water (Fig. 8). Although the beaker was dry on the outside at the beginning of the experiment, it will soon con- dense drops of water from the flame. The water may be easily collected and identified. In burning, hydrogen takes oxygen from the air ; in fact, every case of combustion in air consists in the union of the substance burned with the oxygen of the air. Hydrogen does not, however, support combustion ; that is to say, burning wood, paper, illumina- ting gas, etc., the ordinary combustibles, do not continue to burn when placed in an atmosphere of hydrogen. FIG. s. Hydrogen is really inert at ordinary tem- peratures, and active only when the temper- ature is raised considerably. This is shown by the fact that even hydrogen and oxygen may be mixed and kept together for an indefinite period without 16 HYDROGEN. evidence of action. The inertness of hydrogen under ordinary conditions is such that it is often conven- ient to experiment with a substance in an atmosphere of this gas, thus excluding the active oxygen of the air. At about 350 C., however, hydrogen and oxygen unite with great violence. Although only a very 8mall amount of the mixture of the gases is really heated to 850 C. by the match or electric spark used to start the combustion, yet the burning of this portion causes so much heat to be given off that adjacent portions are quickly raised to the required temperature ; as a result, union takes place very rapidly through the whole mixture. The explosion seems instanta- neous, but is not ; its rate has been determined by photog- raphy to be a little less than two miles a second. If the union of oxygen and hydrogen takes place at a pressure lower than the ordinary atmospheric pressure, the rate of the explosion and, consequently, its violence, is much diminished. Hydrogen has the power of combining not only with free oxygen, but also, in many cases, with oxygen that is in combination with other elements. Thus, if hydro- gen is passed over heated copper oxide and lead oxide, it unites with the oxygen of these substances, forming water, and setting free copper and lead, respectively. These are illustrations of the reducing, i. e., deoxidizing, power of hydrogen. 13. Occlusion of Hydrogen. A remarkable prop- erty of hydrogen is its absorption in large quantities by certain metals, e. #., platinum and palladium. The PHYSICAL PROPERTIES. 17 hydrogen is said to be " occluded " by these metals. The phenomenon may be illustrated by holding a piece of platinum sponge in a jet of dry hydrogen ; so much heat is evolved during the occlusion of the hydrogen that the latter is set on fire. The ready union of hydrogen and other inflammable gases with air, in the presence of platinum, may be shown by directing the mixture of cold gases issuing from a Bunsen burner against a piece of hot platinum foil. The gases will unite with so great an evolution of heat that the platinum will con- tinue to glow. Palladium has an even greater power of occluding hydrogen than platinum has. One volume of palladium can absorb about 375 volumes of hydrogen at the ordinary temperature I 14. Physical Properties. As we have already learned, hydrogen is a colorless and odorless gas. It is the lightest substance known, air being 14 and oxy- gen 16 times as heavy. The weight of one liter of hydrogen at C. and 760 mm. pressure is about .09 gram ; there are, therefore, no less than 11 liters of the gas to the gram. Hydrogen is the standard of density in the case of gases ; its relative density is 1. The rate at which hydrogen diffuses, i. e., mixes with other gases, is four times that of oxygen. A special form of diffusion, viz., transpiration, may be illustrated as follows : 18 HYDROGEN. A porous cup (Fig. 9) is attached securely to a glass tube ending under water. If, now, a bell-jar or a large bottle filled with hydrogen is placed over, and enclosing, the porous cup, bubbles of gas will be seen escaping from the lower end of the tube. The explanation of the phenomenon is that the two gaseous media, air and hydrogen, sepa- rated by the porous partition, tend to form a FIG. 9. homogeneous mixture. But the rate at which the hydrogen passes through the partition is so much greater than that of the air, that an increase of volume, and, therefore, of pressure, occurs within the porous cup. Consequently some of the gaseous mix- ture escapes. When the bell-jar is removed, the reverse diffusion takes place. A somewhat simpler form of apparatus is shown in Fig. 10. A wide glass tube has one end cov- ered with a cap of plaster of Paris. If the tube is filled with hydrogen by displacing the air, and the open end is placed at once under water, water will FIG. 10. rise in the tube. Hydrogen is not very soluble in water ; at 14 C. 100 c.c. of water absorb only 1.9 c.c. of the gas. At the ordinary temperature, hydrogen cannot be liquefied by any pressure, however great; but by intense cold, in addition to great pressure, the gas has been con- densed to the liquid and solid state. Liquid hydrogen is a colorless substance less than one tenth as heavy as water. It boils at about 240 C., at the ordinary pressure. CHEMICAL CHANGES. 19 Hydrogen is a better conductor of heat and of elec- tricity than any other gas. 15. Other Methods of Preparation. The action of metals upon acids is only one of many methods by which hydrogen may be prepared ; other general methods are the following : (a) The decomposition of water by the electric current. This operation is called the " electrolysis " of water ; it takes place only when the water contains small quantities of certain substances called " electrolytes.' 1 The electrolyte commonly used is dilute sulphuric acid. (b) The action of certain substances chiefly metals upon water. Some metals, e. is oxygen ; all of this oxygen is given off upon the application of heat. If the potas- sium chlorate is heated by itself, a test tube or retort FIG. 11. (Fig, 11) of hard glass is needed ; for the temperature at which decomposition begins to take place is 350 to 400 C. There are two stages in the decomposition of the potassium chlorate When it is first heated, potassium chloride, contain- ing no oxygen, and potassium perchlorate, containing 46.2% 24 COMMON LABOEATOEY METHOD. 25 oxygen, are formed ; and some oxygen is liberated. For the decomposition of the potassium perchlorate, a much higher temperature is required. The final result is that the potassium chlorate has been broken up completely into potassium chloride and oxygen. Potassium chloride is not volatile at the temper- ature used, and therefore remains behind in the retort. 20. Common Laboratory Method. The evolution of oxygen takes place much more easily, and at about 200 C., if we heat a mixture of potassium chlorate with manganese dioxide, or ferric oxide, instead of potassium chlorate alone. This is the common labora- tory method. Operation. Approximately equal parts by weight of manganese dioxide and potassium chlorate are powdered separately in clean mor- tars, and then mixed carefully on clean, sized paper. Before decomposing the whole mixture, we test its quality by heat- ing a small portion in an open test tube. If there is evidence of violent combustion, or if large sparks appear in the test tube, we reject the mixture and make a fresh one. A few small sparks indicate only traces of impurity (dust, etc.). If the mixture is found sufficiently pure it is placed in a large test tube, or a small flask (they may be of ordinary soft glass), pro- vided with a delivery tube terminating under water. Oxygen comes off readily when a gentle heat is applied to the flask. If the materials used are not of "chemically pure" (c. p.) grade, the oxygen will contain impurities ; these may be re- moved sufficiently by allowing the gas to bubble through sodium hydroxide solution. In the decomposition of potassium chlorate, as in the reactions studied in Chapter I, the relation between the 26 OXYGEN. quantity -of material taken and that of each of the prod- ucts formed is definite and constant. Thus, 122.5 grams of potassium chlorate give 48 grams of oxygen and 74.5 grains potassium chloride. The ease with which potassium chlorate decomposes in the presence of manganese dioxide was unexplained for a long* time ; but it is probably due to the fact that the two substances react to form intermediate com- pounds, which are decomposed again. The heating of a mixture of manganese dioxide and potassium chlorate to about 200 to 250 C. thus results in the decompo- sition of the potassium chlorate, while it leaves the man- ganese dioxide unchanged. 21. Other Methods of Preparing Oxygen. Three other methods of preparing oxygen will be described briefly; these are, (1) Decomposition of mercuric oxide. (2) Decomposition of barium peroxide. (3) Decomposition of manganese dioxide. The first method is the historic one of Priestley, who discovered oxygen in 1771, and of Scheele (pronounced Shala), who discovered it independently in 1774. When mercury is heated in air to a temperature a little below the boiling point of the mercury (mercury boils at 357 C.), it unites .with a definite weight of oxygen to form mercuric oxide, or red oxide of mercury. PHYSICAL PROPERTIES OF OXYGEN. 27 25 parts by weight of mercury and 2 parts of oxygen give 27 parts of mercuric oxide. Mercuric oxide is a stable compound at ordinary tem- eratures ; but when it is heated to a temperature a little higher than that at which it was formed, it is decomposed again into mercury and oxygen. 27 grams of mercuric oxide always yield 25 grams of mercury and 2 grams of oxygen. Barium peroxide, the second substance named above as a source of oxygen, is a white solid which gives up half of its oxygen when heated ; the other half remains in combination with the barium in the compound barium monoxide. Under appropriate conditions barium monoxide takes up from the air as much oxygen as it already holds, and thus forms the per- oxide. Apparatus has been devised in which these changes take place alternately, and large quantities of oxygen are thus produced for sale. 169 grams of barium peroxide give 153 grams of barium monoxide and 16 grams of oxygen. It is evident that the oxygen formed in both the first and the second methods is taken from the air. Manganese dioxide, the third substance named above, is called, also, " black oxide of manganese.' 1 '' It is found in nature as the mineral pyrolusite. Manganese dioxide decomposes at about 600 C., giving off a third of its oxygen. 261 grams of manganese dioxide give 229 grams of manga- nous-manganic oxide and 32 grams of oxygen. 22. Physical Properties of Oxygen. In whatever way it is prepared, oxygen is colorless, odorless, and tasteless, if pure. It is somewhat heavier than air, and 28 OXYGEN. sixteen times as heavy as hydrogen. One liter of oxygen at C. and 760 mm. pressure, weighs 1.43 grams. Gaseous oxygen may be condensed at 118 C. and 50 atmospheres (= 50 X 760 mm.) pressure, to a bright blue liquid. Oxygen is more than twice as soluble in water as hydrogen ; 100 c.c. of water dissolve about 4 c.c. oxygen under ordinary conditions. Oxygen is the most abundant element, composing about half of the earth's solid crust, eight-ninths of the water, and 23% by weight of the atmosphere. It is an essential constituent of all living things. 23. Chemical Properties. The chief chemical prop- erty of oxygen is its energetic support of combustion ; for substances that burn in air burn much more rapidly in oxygen. Thus, a pine splinter which is merely glow- ing in the air will burst into flame if put into oxygen. Combustion in air is more slow than in oxygen, because the oxygen of the air is diluted with almost four times its volume (= more than three times its weight) of inert gases which do not support ordinary burning at all. Examples of substances which burn readily in oxygen are : Iron, which burns with scintillation, forming the magnetic oxide of iron; magnesium, phosphorus, and sulphur, which burn with intensely brilliant flames ; and charcoal, which burns with a yloiv, as in air, but much OXIDES. 29 more brightly. These substances unite with oxygen in the following proportions : - 21 grams of iron and 8 grams of oxygen give 29 grams of magnetic iron oxide. 3 grains of magnesium and 2 grams of oxygen give 5 grams of magnesium oxide. 31 grams of phosphorus and 40 grams of oxygen give 71 grams of phosphorus pentoxide. 1 gram of sulphur and 1 gram of oxygen give 2 grams of sulphur dioxide. 3 grams of carbon (charcoal) and 8 grams of oxygen give 11 grams of carbon dioxide. The magnetic oxide of iron is a black solid ; magne- sium oxide and phosphorus pentoxide are white solids - the latter is very soluble in water ; sulphur dioxide and carbon dioxide are colorless gases. Sulphur dioxide has the characteristic odor of burning sulphur. 24. Oxides. Because of the readiness with which oxygen unites with other elements, the most common compounds in which elementary bodies are found are oxides. Fluorine and the argon family alone, of all the elements, form no oxygen compounds so far as known. } 7 oxides maybe made directly from the elements, e. g., water, and copper oxide ; but many must be made indirectly, e. g., platinum oxide. To obtain the latter substance the metal platinum must first be converted into other compounds, and these into the oxide. Nitrogen, too, although incombustible in the ordinary sense, yet forms, by indirect methods, five dif- ferent oxides. 30 OXYGEN. 25. Oxidation and Reduction. To the union of oxygen with other substances we give the name oxida- tion, and to a substance which gives up some of its oxygen to another body the name oxidizing agent. Of course the oxidizing agent is itself reduced, i. e., loses oxygen, when it oxidizes another substance there can be no oxidation without a corresponding reduction. The substances mentioned in this chapter as sources of oxy- gen are all oxidizing agents ; carbon and hydrogen (c/. 12) are common examples of reducing agents. 26. Deflagration. When a solid or liquid combus- tible substance is mixed with a solid or liquid oxidizing agent, and the temperature is raised sufficiently, the combustion does not proceed from one part of the com^ bustible to another, as is the case when the combustible burns in air; on the contrary, union takes place almost instantaneously through the whole mixture, just as it does through a mixture of gaseous hydrogen and oxy- gen. This rapid union of combustible and oxidizing agent is called deflagration. A common case of deflagration is that of gunpowder, which is a mixture of charcoal and sulphur (reducing agents) with potassium nitrate, or potassium chlorate (oxidizing agents). When ignited in air, gunpowder deflagrates. In an enclosed space the same action takes place, but the gaseous products of the combustion are held for an instant under great pressure. When this pressure is released, the expanding gases are capable of hurling a projectile, or of tearing apart masses of rock. SLOW COMBUSTION. 31 27. Combustion. When a substance unites directly with gaseous oxygen we speak of the oxidation as a case of burning, or combustion. Of this we generally distinguish two kinds, (1) ordinary and (2) slow com- bustion. In ordinary combustion, heat is produced by the union of combustible with oxygen much more rapidly than it can be dispersed by conduction, radiation, etc. ; conse- quently the temperature of the burning body rises far above that of the surrounding medium. Usually a part of the combustible, or of the products of combustion, becomes incandescent ; and some of the energy liberated appears in the form of light. 28. Slow Combustion. A slow combustion occurs when oxidation takes place through a long period of time, and, therefore, without a decided rise of temper- ature. This can occur only when the heat is dispersed as rapidly as it is evolved. Decay, e.g., of wood, is a form of slow combustion ; so is the rusting of iron. The temperature of the bodies of animals is kept up by the slow oxidations taking place within them. Although there is a great difference in temperature between a body oxidizing slowly and one burning in the ordinary way, yet the amount of energy actually evolved by the oxidation of a given weight of a sub- stance is the same in the one case as in the other. -i 32 OXYGEN. Thus, a piece of magnesium oxidizes slowly in moist air, at the ordinary temperature, to form a white powder containing magnesium oxide and water ; but we have every reason to believe that the quantity of energy set free during this slow formation of magnesium oxide is just as great as that evolved when an equal mass of the metal magnesium burns brightly in the air. That there is no perceptible heat and light in the former case is due to the fact that the evolution of heat is equaled by its dispersion. 29. Spontaneous Combustion. When the evolu- tion of heat is only a little in excess of its dispersion, the combustion is apparently a slow one ; after a while, however, enough heat accumulates to set the body on fire. Slow oxidation explains so-called- " spontaneous " combustions, by which heaps of oily rags, etc., ignite without apparent cause. Spontaneous combustion may be illustrated by means of a solution of phosphorus (only a small amount must be used) in carbon disulphide. If this solution is poured upon a tilter paper supported on a ring stand, the phosphorus will soon take fire "spontaneously." The explanation of the phenomenon is that the evaporation of the carbon disulphide leaves the phosphorus in the pores of the paper, where it oxidizes ; and the heat generated, being prevented from escaping by the non- conducting filter paper, soon raises the temperature of some part of the paper to the ignition temperature of the phos- phorus, 30. Ignition Temperature. In all ordinary com- bustion there is a definite temperature, called the ignition temperature, or the kindling temperature, to COMBUSTION IN AIR; DRAFTS. 33 which the combustible substance must be heated in order that it may begin to unite with the gas supporting the combustion. The burning substance must not only be heated up to the kindling temperature, but it must be kept at least as liigli as this temperature, or combus- tion will cease. The ignition temperature is different for different sub- stances. Thus, ordinary phosphorus bursts into flame, in air, at about 40 C. ; while for sulphur the kindling tempera- ture is about 260 C. In some cases the temperature of ignition is far below the ordinary temperature. The heat evolved by the burning of one part of a substance serves to raise other parts to the ignition point; illustrations of this are found in the burning of wood, paper, etc. In the ordinary match the ignition temperature of the material composing the head is reached by friction; and the heat generated by the combustion of the head serves to raise the wood of the match to its ignition temperature. 31. Combustion in Air; Drafts. The heat given off in combustion is taken up, not only by new portions of the burning body, but also by fresh portions of the surrounding gaseous medium ; if, therefore, the latter is diluted, some of the heat available will be used in raising the temperature of the diluting gas, as well' as that of the gas taking part in the combustion. Hence, in the air, which is a mixture of oxygen and nitrogen 34 OXYGEN. (chiefly), combustion is much slower than in pure oxy- gen, since the nitrogen, although it takes no part in the combustion, yet takes up much of the heat evolved by the combustion. Moreover, when the products of the combustion are gaseous, they dilute the oxygen still further. Thus, when charcoal burns in an enclosed portion of air, com- bustion ceases long before all the oxygen is exhausted, for the reason that the carbon dioxide gas which is formed dilutes the oxygen. If the products of the combustion are removed as rapidly as formed, the combustion will be much more nearly com- plete. Thus, when phosphorus burns in an enclosed portion of air over water, the product of the combustion phosphorus pentoxide dissolves in the water, and does not dilute the oxygen ; as a result the oxygen is practically all taken up by the phosphorus. For the same reason the combustion of charcoal in an enclosed portion of air may be made much more nearly complete if carried out over sodium hydroxide solution. The removal of the products of combustion from the " sphere of action " is accomplished in ordinary burning by means of drafts, which also bring fresh supplies of air. A moderate draft is thus beneficial to combustion. The air current may, however, have such a velocity that the heat evolved in the combustion is not sufficient to raise the temperature of the air supplied and of fresh portions of the burning body to the kindling tempera- FLAMES. 35 cure. Hence combustion ceases. A flame may thus be " blown out " by a strong current of air. 32. The Safety Lamp. The lowering of the tem- perature of a flame below the kindling temperature is admirably illustrated in the safety lamp devised by Sir Humphry Davy. The lamp consists of an ordinary lantern entirely surrounded by wire gauze. When such a lamp (Fig. 12) is carried into an explosive mixture of gases, e.g., hydrogen and air, the gases diffuse through the wire gauze and burn inside the lamp ; but the heat generated is conducted away by the wire gauze instead of being communicated to the explosive mix- ture outside ; hence an explosion of the gases outside of the lamp is avoided. Safety lamps are used to prevent the explosion of the " fire-damp," a mix- ture of marsh gas and air, which often occurs in mines. FIG. 12. 33. Flames. Aflame is a gas in combustion. To burn with a flame, a substance must either be gaseous itself, or it must evolve gaseous products. Such sub- stances as magnesium, sulphur, phosphorus, and wax 36 OXYGEN. burn with flame because they are first converted into the gaseous form; wood and soft coal, because they evolve combustible gaseous products ; but charcoal, which contains practically 110 volatile constituents, merely glows. The structure of flames will be taken up later. 34. Reversed Combustion. We have spoken of combustion heretofore as the union of the burning body with oxygen ; other gases, however, may be supporters of combustion just as truly as oxygen. Thus, a jet of burning hydrogen continues to burn in bro- mine vapor ; and phosphorus burns in chlorine much as in oxygen. If both the combustible and the supporter of its coin, bustion are gaseous, the combustion may be reversed. Thus, oxygen may become the burning body and illumi- nating gas the supporter of combustion. This reversal may be shown by a very simple experiment. A bottle (Fig. 13) is supported, mouth do\vn- ward, and filled with illuminating gas by dis- jL- /xC placing the air. The gas at the mouth of the bottle is then lighted, and while it is burning ^ """ a jet of oxygen is brought up into the bottle. FIG. 13. The oxygen takes fire at the bottle's mouth and burns in the atmosphere of illuminat- ing gas. The oxygen jet of the preceding experiment may EXERCISES. 37 be replaced by a deflagrating spoon of potassium chlorate which has been heated so that it gives off oxygen. 35. Exercises. 1. How many grams of mercury will b^ formed by the de j composition of 43.2 grams of mercuric oxide ? How many grams of oxygen ? 2. How many grams of the magnetic oxide of iron will be formed when 50 grams of iron burn in oxygen ? 3. How many grams of manganese dioxide are needed to give, when decomposed by heat, 12 grams of oxygen? How much manganous-manganic oxide is formed at the same time ? 4. How much magnesium is contained in 30 grams of mag- nesium oxide ? What per cent of magnesium oxide is mag- nesium ? Oxygen ? 5. Calculate the per cent of oxygen in phosphorus pen- toxide. 6. How many c.c. of oxygen can be made from 1.2 grams of potassium chlorate when 1 c.c. of the gas weighs 0.0014 grams ? 7. How many grams of potassium chlorate must be decom- posed to fill a 36-liter gasometer (a vessel for storing gases) with- oxygen at a temperature and a pressure at which 1 liter of oxygen weighs 1.25 grams ? 8. What is the weight of the carbon dioxide formed by the combustion of 10 grams of carbon in oxygen ? What will be the volume of the carbon dioxide under conditions at which one liter of the gas weighs 2 grams ? CHAPTER III. WATER. 36. Nature of Water. The union of the elements composing water is so strong that water itself was be- lie vecl to be an element until 1781. In that year, Cav- endish, who had discovered hydrogen in 1766, succeeded in synthesizing (= putting together) water from hydro- gen and oxygen, and thus proved its compound nature. 37. Electrolysis o f Water. The fact that water is a compound is proved analytically by its decomposition b y the electric current. T h e operation is carried 27.129 g. oxygen j 27.129 : 3.39 :: x : 1 ; whence x = 8.002 -f , the ratio of oxygen to hydrogen in water. 40. Natural Water and Its Impurities. The water that falls upon the earth's land surface gets back to the 44 WATER. sea in various ways, but rarely without leaching out soluble substances from the soil. Natural water there- fore contains more or less impurity. The character of the impurity depends, (1) upon the substances present in the air through which the water fell to the earth ; (2) upon the soil through or over which the water has flowed ; and, also, (3) upon the opportunities the water has had of losing material previously gathered. Even rain water is far from pure, for it gathers much dust, both organic and inorganic, and many gaseous impurities of the air, e. g., ammonia. Water which penetrates the earth's surface usually finds soluble substances, both solid and gaseous. The most common soluble solids found in water are, prob- ably, common salt, magnesium chloride, and gypsum; of the gases, carbon dioxide and hydrogen sulphide. Water charged with carbon dioxide has the power to dissolve limestone ; hence this substance is a common ingredient of natural water, even of moderately "soft" water, as is proved by the incrustations of limestone in vessels in which such water is habitually heated. Water charged with hydrogen sulphide is called sul- phur water. The water which flows over the earth as rivers gathers its peculiar organic impurities from the land. These may consist of micro-organisms washed down by surface water, or, in the case of rivers passing large cities, of sewage. If the river is sluggish, these impurities are not easily removed ; THE PURIFICATION OF WATER. 45 but if it has a rapid current, and especially if there are rapids and waterfalls in its course, the river soon purifies itself by bringing its impurities into contact with the oxygen of the air, which destroys them. 41 . Sea Water. Since the sea is the ultimate desti- nation of most of the water that falls upon the land, it is evident that the material dissolved by fresh water will accumulate in the ocean. Indeed, about four per cent of sea water consists of dissolved material, three- fourths of which is common salt. It is probable that greater or smaller amounts of all the substances com- posing the crust of the earth may be found in the sea. 42. The Purification of Water. Water may usually be purified by filtration or by distillation. Filtration serves not only to remove insoluble sub- stances, but also to oxidize many organic impurities by bringing them into intimate contact with air. When water is raised to the boiling temperature, most of the micro-organisms contained in the water are killed, and at least one of its inorganic impurities, viz., calcium carbonate (lime- stone), is rendered insoluble. Soluble impurities, however, still remain. To get Avater free from these it must be distilled. Distillation consists in converting a liquid into vapor, and then condensing the vapor to the liquid state. When water is distilled, all impurities more volatile than the water will appear in the first portions of the 46 WATER. distillate ; all much less volatile will remain behind in the retort. The usual form of distilling apparatus used in labora- tories is shown in Fig. 18. The condenser in the figure is called a Liebig's condenser. It consists of an inner tube through which the evolved vapor is passed for condensation, and of an outer jacket through which a stream of cold water is kept running in the direction shown by the arrows. FIG. 18. Although water can be obtained reasonably pure by distillation, chemically pure water is very difficult to prepare. Even if water is pure when freshly distilled, as shown by its leaving no residue when evaporated in a platinum dish, it cannot be kept pure long, owing to PROPERTIES OF WATER. 47 its tendency to act upon the glass or porcelain vessels in which it is stored or used. Distilled water is " flat " to the taste. This is due largely to the fact that distilled water has lost the gases present in natural water. Distilled water may therefore be made much more palatable by shaking it thoroughly with air. 43. Hard and Soft Water. Water which contains much gypsum, limestone, or similar substances in solu- tion does not wet the skin readily, and is therefore called "hard" water. When soap is put into such water it does not dissolve readily, but forms an insoluble scum. It is only after the separation of this scum that soap will dissolve in quantity and form permanent suds. As will be explained later, the hardness of water containing only limestone is temporary because it may be removed by boiling ; if gypsum is present, however, the water is permanently hard and can be " softened " only by the use of washing powders, etc., which are capable of convert- ing the gypsum into insoluble forms. 44. Properties of Water. Pure water is practically odorless and tasteless. In small quantities it has no color, but in large masses it is blue. The specific heat of water is high, more heat being required to raise the temperature of a given weight of water one degree than is required in the case of any other substance except hydrogen. 48 WATEtt. The latent heat both of water and of steam is very great. When a given weight of water at C. is frozen to ice at C., it gives off enough heat to raise the temperature of an equal weight of water from C. to 80 C. When a given weight of steam at 100 C. con- denses to water at 100 C., the heat evolved is sufficient to raise the temperature of about 5.37 times the weight of water from C. to 100 C. The boiling point of water is 100 C. at 760 mm. pressure. Since the boiling point of a substance is the temperature at which the pressure of its vapor just ex- ceeds that of the atmosphere, the " vapor tension " of water at 100 C. must be 760 mm. At 21 mm. pres- sure, water boils at 23 C. ; at 3,581 mm., at 150 C. The freezing point of water (= melting point of ice) is C. at 760 mm. pressure. Water expands on freez- ing, 10 c.c. of the liquid becoming about 10.1 c.c. of ice. At about 4 C., water is at its maximum den- sity. Its relative density at this temperature is taken as 1. Water is a poor conductor of heat and of the electric current. 45. Steam and its Dissociation. Steam, which is water in the condition of a vapor, is nine times as heavy as hydrogen. It is so stable that it does not begin to decompose into its elements until it is heated to about 1000 C. Above 1000 C. the amount of decomposition increases with the temperature, until at 2500 C. about ACTION OF SODIUM UPON WATER. 49 half of the steam is no longer steam, but oxygen and hydrogen uncombined. No matter how long steam is kept at 2500 C. it cannot be decomposed completely, for the reason that side by side with the decomposition of steam into hydrogen and oxygen there is a recombina- tion of these elements to form steam. At every tem- perature, therefore, between 1000 C. and the (high) temperature at which the decomposition of steam is complete, a condition of equilibrium is soon reached, at which as much steam is produced in a given time as is decomposed in the same time. Hence the change pro- ceeds no farther unless the temperature is raised. If the temperature is lowered, enough hydrogen and oxy- gen recombine to produce equilibrium at the lower temperature. A decomposition like that of steam is called a disso- ciation. 46. Action of Sodium upon Water. In the first chapter reference was made to the action of sodium upon water as a means of preparing hydrogen. This reaction will now be considered more fully. Sodium is a soft solid, somewhat lighter than water. It has a silvery luster when freshly cut, but tarnishes quickly in ordinary air. It is, therefore, kept under lig- roin or kerosene. Although we commonly think of a metal as hard and heavy, sodium, which has neither of these properties, is yet one of the best representatives of the class of metals. This is due to its chemical prop- 50 WATER. erties, the most important of which, for our present pur- pose, is its behavior toward water. When a piece of sodium is thrown upon water (this is done at arm's length to avoid danger from spattering), it at once attacks the water, melts, assumes a globular form, and then swims about until dissolved. If a lighted match is held near the sodium while it is floating upon water, a flame will appear ; this is burning hydrogen. The hydrogen may be collected by placing the sodium in a short piece (1 cm. long) of glass tubing, and then plunging it quickly by the aid of tongs under the mouth of a bottle, or a test tube, filled with water and inverted in a pan of water. Water in which a sufficient quantity of sodium has dissolved possesses new properties. It feels soapy to the touch, has a bitter taste, and turns many vegetable col- ors, e. g., red litmus to blue. If the water is evaporated, a white substance will remain, which is 'sodium hydrox- ide, or caustic soda. It is this substance that gives the water its new properties. 47. Quantitative Study of the Reaction ; Hydrox- ides. Sodium hydroxide is, as its name indicates, a compound of sodium, hydrogen, and oxygen. Since sodium hydroxide contains hydrogen, it is evident that not all of the hydrogen of water is set free by sodium. As a matter of fact, only one-half of the hydrogen is so liberated, the remaining half being in combination with sodium and oxygen. This appears from a quantitative study of the reaction. STUDY OF THE REACTION; HYDROXIDES. 51 If 23 grams of sodium had been placed in contact with a quantity of water greater than 18 grams, it would have acted upon 18 grams of water only, and would have formed 40 grams of sodium hydroxide, and, at standard temperature and pressure, a little over 11 liters of hydrogen. The hydrogen formed would weigh prac- tically 1 gram. To summarize the results of the reac- tion, quantitatively as well as qualitatively : 23 grams of sodium and 18 grams of water react to give 40 grams of sodium hydroxide and 1 gram of hydrogen. As we have already learned, 18 grams of water con- sist of 2 grams of hydrogen and 16 of oxygen. Further- more, 40 grams of sodium hydroxide would give, when decomposed, 23 grams of sodium, 16 grams of oxygen, and 1 gram of hydrogen. It is evident, therefore, that sodium replaces only half of the hydrogen of water in forming sodium hydroxide. If we had used the metal potassium, the results would have been similar, viz. : 39 grams of potassium and 18 grams of water react to pro- duce 56 grams of potassium hydroxide and 1 gram of hydrogen. The 56 grams of potassium hydroxide consist of 39 grams of potassium, 16 grams of oxygen, and 1 gram of hydrogen. Here again one-half of the hydrogen of 52 WATER. the water decomposed is liberated ; and the remaining half is retained in the hydroxide. In fact, the hydrox- ides of all of the metals may be considered to be water with half of its hydrogen replaced by a metal. 48. The Action of Metals upon Hydroxides. By the use of proper methods, the hydrogen of sodium hydroxide may be replaced by sodium, and that of potassium hydroxide by potassium. The resulting sub- stances will be sodium and potassium oxides. The re- actions take place as follows : h 40 grams of sodium hydroxide and 23 grams of sodium give 62 grams of sodium oxide and 1 gram of hydrogen. Also, 56 grams of potassium hydroxide and 39 grams of potassium give 94 grams of potassium oxide and 1 gram of hydrogen. To replace the hydrogen of sodium hydroxide by sodium, and that of potassium hydroxide by potassium is a somewhat difficult operation; but it is very easy to replace the hydrogen of these hydroxides by alumi- num. The resulting substances are sodium-aluminum oxide and potassium-aluminum oxide * respectively. These compounds are similar to sodium oxide and to potas- sium oxide in that they are water with all of its hydro- gen replaced in two stages by metallic elements. All of the facts just stated are given in the following recapitulation : * Cf . aluminates, 424. WATER MECHANICALLY ENCLOSED. 53 HYDROGEN. OXYGEN. SODIUM. POTASSIUM. 18 parts by weight of water consist of 2 16 23 40 parts by weight of sodium hydroxide consist of 1 16 62 parts by weight of sodium oxide consist of 1 16 16 16 46 56 parts by weight of potassium hydroxide consist of 39 94 parts by weight of potassium oxide consist of 78 49. Water in Combination. Water is widely dis- tributed, not only in the free condition, but also in combined form. Most natural substances contain it. This is true not only of animal and plant tissues and products, but even of inorganic substances. For con- venience, 1 we may distinguish at least three ways in which water may be contained in other substances : (1) Mechanically enclosed. (2) As " water of crystallization." (3) As an integral part of the substance. The form in which water is contained in a substance gener- ally appears from the behavior of the substance when heated. 50. Water Mechanically Enclosed. Water may be held mechanically either (1) between the crystals of a substance, or (2) in its pores. In either case the water 54 WATEE. is given off when the substance is heated gently. When the substance contains water enclosed between crystals, however, the water escapes explosively ; for the crystal-mass is broken in pieces by the steam pro- duced. Such substances are said to decrepitate. Illustrations are : Common salt and potassium sulphate. 51. Water of Crystallization. By water of crys- tallization^ or crystal-water, we mean the water with which some substances combine when they crystallize from aqueous solution. When substances containing water of crystallization are heated, they usually melt while the water escapes, and then assume the solid form again. The loss of crystal-water by a substance is accompanied by a loss of crystalline structure and by other changes in proper- ties. Thus, cupric sulphate is a white solid, but blue vitriol, its ordinary form, is cupric sulphate plus water of crystallization. Crystallized sodium sulphate, sodium carbonate, alum, etc., all contain much crystal-water. The amount of crystal-water which Avill combine with a given weight of the anhydrous (i.e., water-free) sub- stance is definite for each substance. Thus, 90 grams of water are united with 159 grams of cupric sulphate in every 249 grams of blue vitriol. The color of a substance is usually the same when EFFLORESCENCE, DELIQUESCENCE, ETC. 55 the substance is combined with water of crystallization as when it is in solution in water. By no means all crystalline substances contain crystal- water. Cane sugar, salt, potassium sulphate, etc., crystallize from their solutions in water without taking up any of the water as crystal-water. 52. Water an Integral Part of the Substance. Many substances which cannot be said to contain water, yet contain hydrogen and oxygen in the proportion in which these elements are united in water. Such compounds are the substances referred to as having water in the third form of combination, viz., as an integral part of the substance. Examples are sugar, starch, cotton, wood, etc. When these substances are heated, they are easily decomposed, liberating water and other products, while a residue of charcoal remains behind. 53. Efflorescence, Deliquescence, Etc. Certain substances give up all or part of their crystal-water when exposed to the air, and thus lose their crystalline form. Such substances are said to effloresce. An example is crystallized sodium carbonate, which be- comes a non-crystalline powder when exposed to the air. Certain other substances, on the contrary, when de- prived of their water of crystallization, take it up 56 WATER. again by absorbing water from the air and from other substances. Such bodies make good dehydrating, i. e., drying, agents. Examples are: Anhydrous cupric sulphate and anhydrous potassium carbonate." If dehydrating agents absorb so much water from the air that they dissolve in the water, they are said to deliquesce. An example is anhydrous calcium chloride. In any case, if a substance takes up water when ex- posed to moist air, it is said to be hygroscopic. Many substances are drying agents, not because they tend to take up water of crystallization, but because they combine with water to form other compounds. Quicklime, which is calcium oxide, is an example. When this sub- stance is slaked, by addition of water, it becomes calcium hydroxide. f 54. Exercises. I 1. How "many grams of water are formed by the combustion of 10 grams of hydrogen in air ? 2. What evidence is there that the hydrogen of water is more divisible than the oxygen ? 3. How could you determine approximately how much water is contained in a potato ? EXERCISES. 57 4. 5 grams of crystalline barium chloride lost, when heated at 120 C., 0.65 grams. What per cent of water did it have ? 5. Calculate the per cent of water of crystallization in a sample of potash alum, 47.4 grams of which lost. 21. 6 grams of water. 6. How many grams of water can be decomposed by 10 grams of sodium? How much sodium hydroxide will be formed ? 7. How many grams of potassium are required to give with water 50 grams of potassium hydroxide? How many grams of hydrogen will be liberated at the same time ? How many liters when 1 liter weighs 0.09 grams? 8. How many grams of water are formed by burning 10 liters of hydrogen when 1 liter of the latter weighs 0.085 grams ? CHAPTER IV. SOLUTION. 55. The Character of Solution. Solution, or dis- solving, takes place when substances are mixed in such a way that the matter of each is distributed uniformly through that of the others. The resulting homogeneous mixture is called a solution. Thus broadly denned, the term solution includes phenomena called by many dif- ferent names, but we shall restrict it to the absorption of a gas, liquid, or solid, within the portion of space occupied by some liquid. The liquid is called the solvent. Examples of common solvents are : Water, alcohol, and ether. If the solvent is colorless, and the dissolved substance has a definite color, the solution will usually be colored ; if the dis- solved substance is white, or colorless,. the solution will be colorless ; but in any case the solution will be clear. Insoluble substances often remain mechanically suspended in a liquid for some time. Their presence is shown by the turbid appearance of the, liquid. Dilute solutions have practically the same volume as that of the solvent, but when the amount of dissolved substance becomes large, the volume of the solution is increased. 68 TEMPERATURE CHANGES. 59 An illustration of the first statement is found in the familiar experiment in which a considerable quantity of powdered sugar is added, little by little, to a vessel entirely full of water without causing an overflow, while a much smaller amount of an insoluble substance, e. g., sand, causes some of the water to be displaced. 56. Boiling Point and Freezing Point of a Solution. Substances in solution raise the boiling point, and lower the freezing point of the solvent. Thus, water con- taining salt or sugar boils above 100 C. at 760 mm. pressure, and freezes below C. In dilute solutions the rise of the boiling point and the lowering of the freezing point are proportional to the amount* of dissolved substance in a given volume. The specific gravity of solutions of solids is greater than that of the solvent. Thus, sea water has a specific gravity of 1.026. 57. Temperature Changes during Solution. When a gas dissolves in a liquid there is, as a rule, an evolu- tion of heat and a consequent rise of temperature, but the solution of a solid is usually attended by an absorp- tion of heat and a reduction of temperature. Some solids, however, dissolve in water with evolution of heat. Examples are : Anhydrous calcium chloride and anhydrous sodium carbonate. These apparent excep- tions are usually substances that have been deprived of water of crystallization, and take it up again when brought into contact with water. Because of the heat 60 SOLUTION. evolution due to the union of these substances with their crystal-water, the heat absorption due to the solution of the crystallized substances is not perceptible. When the crystals of such substances are dissolved in water there is usually a reduction of temperature. 58. Solubility. By the solubility of a substance we mean the maximum amount of the substance that can be taken up by a given quantity of the solvent under the given conditions. The amounts of two substances which will dissolve in a given weight of a solvent are very unequal, as are, also, the quantities of two solvents which are required to absorb a given weight of the same substance. Thus, sugar and salt are very soluble in water, but practically insoluble in ether. Even in water, however, their solubilities are very different, sugar being much more soluble than salt. Similar differences exist in the case of gases, hydrogen, for example, being only about half as soluble as oxygen in water of the ordinary temperature. 59. Effect of Temperature on Solution. -- The solu- bility of a substance depends not only upon the solvent used, but also upon the temperature. As a rule, solids are more soluble in hot than in cold liquids, while the reverse is true of gases. The following table shows the effect of temperature upon the solubility of several solids, SOLUBLE AND INSOLUBLE SUBSTANCES. 61 SUBSTANCE. GRAMS SOLUBLE IN 100 GRAMS WATER. AtOC. At 20. At 100. Potassium nitrate. 13.3 31.7 246. Sodium chloride. 35. " 36. 39.7 Potassium chlorate. 7.2 59.5 Cupric sulphate (cryst.). 42.3 203.3 To illustrate the decrease of solubility of gases with rise of temperature we may take the case of oxygen, 4.1 c.c. of which can dissolve in 100 c.c. of water at C., 2.9 c.c. at 15, and none at 100. 60. Soluble and Insoluble Substances. A solid requiring less than 100 times its weight for com- plete solution may be considered readily soluble ; one needing between 100 and 1,000 times its weight, difficultly soluble ; while one which requires more than 1,000 parts of the solvent may be called insoluble. There is, however, a great difference in the solubilities of so-called " insoluble " substances. Thus, strontium sulphate requires about 8,000 parts of water for solution, and barium sulphate, about 400,000 parts. We call a substance insoluble, then, only relatively to other substances, and not absolutely, for there is prob- 62 SOLUTION. ably no substance of which a small amount will not dissolve, if the quantity of the solvent is very large. 61. Saturated Solution. When a liquid has in solu- tion all that it can hold of a substance under certain conditions, it is said to be saturated with respect to that substance under the specified conditions. Considering now only the solution of solids in liquids, we may know that a solution is saturated when a slight lowering of its temperature or the removal of a small amount of the solvent, e, g., by evaporation, causes precipitation of some of the dissolved substance. There are two methods in common use for the production, at ordinary temperatures, of a saturated solution of a solid. In the first of these the solvent is allowed to remain for some time in contact with an excess of the solid, and the mixture is shaken or stirred to hasten solution. In the second meth'od, the solvent is heated above the ordi- nary temperature with enough solid to produce saturation at the higher temperature, and the solution is then cooled to the ordinary temperature. In this way, the excess of solid is deposited. 62. Supersaturated Solutions. -- Many solutions, however, although saturated at temperatures above the ordinary, will not deposit their excess of dissolved solid when the temperature is lowered. Such solutions are said to be supersaturated. A solution usually remains supersaturated only while undisturbed. If the contain- ing vessel is jarred, or if small particles, e. #., of dust, P&ECIPiTATION AND CttYSTALLlZAflOfr. 63 are introduced, precipitation often results. The most certain way, however, of disturbing a supersaturated so- lution is to add a crystal of the dissolved substance. A rapid separation of the excess of the latter is the result. 63. Precipitation and Crystallization. As stated in the preceding section, a solid may be made to separate from its solution, if the latter is brought to the point of saturation. A dilute solution must, therefore, be con- centrated if separation is to take place. If a- solid sep- arates from solution rather slowly, it will frequently be found to consist of regular masses called crystals. The more sloivly crystallization takes place, the larger and more perfect will the crystals be. But often a solution is brought to saturation suddenly, as is the case when the temperature of an almost saturated solution is rapidly lowered, or when another solvent is added, or when a new substance is formed which is not very soluble in the solvent. In such cases, the substance that separates from solution will consist of very small crystals, or it may even be in an amorphous, i. e., non-crystalline, form. In either case it is called a precipitate. Thus, when silver nitrate and sodium chloride solutions are mixed, there is produced a white, amorphous precipitate of silver chloride, the sodium nitrate formed at the same time remaining in solution. Relatively to silver nitrate solution, therefore, sodium chloride solution is a precipitant, since its addition produces a precipitate. 64 SOLUTION. 64. Decantation and Filtration. We may separate a precipitate from the solution in which it is suspended either by allowing it to settle and then decanting, i. e., pouring off, the clear solution, or by filtering the mixture of liquid and solid. For the latter purpose we use a filter paper, consisting of cellulose, which permits liquids and dissolved solids to pass through its pores, but usu- ally holds back suspended solids. What passes through the filter is called the filtrate. 65. Crystallization from Fusion. A solid may sep- arate in crystalline- form not only from solution, but also by the solidification of a liquid, i. e., from fusion. Thus, water freezes, and molten sulphur solidifies, in crystal- line form. Just as there is a supersaturated condition of some so- lutions, owing to a tardy separation of dissolved solids, so there is a super/used condition of some liquids, because of their slow assumption of the solid form even at temperatures below their true freezing points. Crys- tallization is effected in the same way in both cases, viz., by " inoculation " with a crystal of the solid. 66. Effervescence. Gases, like solids and liquids, separate from solution if formed in a solvent unable to hold them. Because of their low specific gravity, how- ever, gases rise to the top of the solution, and thus escape into the air. A liquid evolving a gas is said to effervesce. EXERCISES. 65 The action of zinc on dilute sulphuric acid, for example, causes effervescence of the dilute acid, owing to the escape of hydrogen. Similarly, "soda water" effervesces, because of liberation of carbon dioxide gas. 67. Exercises. 1. Why do subterranean waters contain more gaseous sub- stances in solution than surface waters ? 2. Why is it that insoluble substances, e. , or =). The resulting ex- pression is called a formula. Thus, mercuric oxide, a compound of mercury and oxygen, is represented by the formula HgO. Similarly, HC1 represents a compound of hydrogen and chlorine, i. e., hydrochloric acid. 71. Symbolic Equations. A symbolic equation (called, simply, " an equation ") is formed by writing, instead of the names of the reacting substances, their SYMBOLIC EQUATIONS. 69 symbols and formulas. Thus, the fact that sodium acts upon water to produce sodium hydroxide and hydrogen is shown by the expression, Na + HOH (or H 2 O) NaOH + H. In this equation Na is the symbol of sodium ; H, that of hydrogen ; and O, that of oxygen. The formula HOH, or H 2 O, for water, shows that water is composed of hydrogen and oxygen, while the formula NaOH, for sodium hydroxide, shows that sodium hydroxide is a compound of sodium, oxygen, and hydrogen. The sign (or =) is best read "give," or "produce." The sign -f- is read "and." The symbols and formulas of the equation are read in the direction of the arrow ; those preceding the arroAV are called the factors, those succeeding it the products of the reaction. When the sign = is used, the symbols and formulas to the left of it are the factors, and those to the right the products. Among the advantages of the use of symbols are the fol- lowing : 1. Symbolic expressions for compounds enable us to see, at once, what elements make up the compounds. Thus, while the common names of salt, water, galena, and caustic soda give us no idea of the composition of these bodies,- the corresponding formulas, NaCl, II 2 O, PbS, and NaOH, do. 2. Symbolic equations, if correctly written, enable us to de- termine what changes in the relations of elements have resulted from a given case of chemical action. Thus it requires only a glance at the symbolic equation for the action of sodium upon 70 FUNDAMENTAL LAWS. water to determine that in this reaction sodium displaced part of the hydrogen of water to form sodium hydroxide. 72. Equations the Result of Experiment. Equa- tions mean nothing unless they are the result of experiment. If the student cannot himself prove what products are formed in a reaction, he must depend upon some trust- worthy source, e. g., the teacher or a text-book, for the necessary information. The student should learn the important equations given from time to time in the text, hot by rote, but with appreciation of their mean- ing. He will soon find that the number of equations that must actually be memorized is small, for, with a few typical equations as a basis, a large number of analogous ones can readily be acquired. 73. Quantitative Meaning of Symbols and Equa- tions. The equation Na + H 2 O NaOH + H means to the chemist much more than has been stated in 10, for it indicates the proportions by weight of the reacting substances and of the products. We can give this added meaning to the above and to every equation by letting the symbol of each element represent not only the element in general, but also a definite mass of it. The formula of a compound will then show not only what elements are contained in the compound, but, in addition, the proportions by iveight in which they are united. Finally, if we assume that symbols and forniu- SYMBOLS AND FORMULAS. 71 las stand for definite masses of elements and of com- pounds, respectively, then the equation for every reaction will represent the proportion by weight of every substance entering into the reaction. Thus, the equation Ka + H 2 O NaOH + H means to the chemist that 23 parts by weight of sodium react with 18 parts of water to produce 40 parts of sodium hydroxide and 1 part of hydrogen. Of course the proportions will be true, whatever units are used, i. e., with pounds and tons as well as with milligrams and grams, but for our present purpose we will let symbols and formulas represent grams. Na means, therefore, 23 grams of sodium, and H 2 O, 18 grams of water. Since water is one-ninth hydrogen and eight-ninths oxygen, 18 grams of water must contain 2 grams of hydrogen and 16 of oxygen. O, therefore, means 16 grams of oxygen. 74. How to Represent Multiples of Symbols and Formulas. In the formula H 2 O, H 2 means twice the quantity of hydrogen represented by H, for in all chemi- cal formulas a small figure written after and slightly below a symbol multiplies the quantity represented by the symbol immediately preceding. The formula NaOH means 40 grams of sodium hydroxide. Of the 40 grams, 23 are sodium, 16 are oxygen, and 1 is hydrogen. The equation thus ac- counts for every gram of material taken: Na + H 2 O NaOH + H. 23 + (2 +16) (23 +16 +1) +1 72 FUNDAMENTAL LAWS. If we multiply the quantities taken in the above equation by 2, the relative amounts are not altered. Thus the equatior 2 Ka-f 2 H 2 O 2 KaOH + H 2 represents all the facts as well, at least, as the simpler equation. A figure written before a formula multiplies the quantity of the substance represented by the formula, just as the small figure written after a symbol multiplies the quantity indicated by the symbol. Therefore 2 ]S T a means 46 (= 23 X 2) grams of sodium ; 2 H 2 O means 36 (=2 [2 + 16] ) grams of water ; 2 XaOH means 80 (=2 [23 -j- 16-j-l] ) grams of sodium hydroxide ; and H 2 means 2 grams of hydrogen. It often happens that a formula contains a group of elements repeated two or more times. Thus, the for- mula of calcium hydroxide is Ca(OH) 2 ; it may also OH be written CaO 2 H 2 ; or, better still, Ca^,j. In the first of these formulas, viz., Ca(OH) 2 , the small figure written after the parenthesis multiplies what is in the par- enthesis immediately preceding, just as if the symbols in the parenthesis were together one symbol. Symbols which are grouped together in this way are called radicals. Other examples of formulas containing radicals are : Cu- (XO 3 ) 2 for cupric nitrate, and Fe 2 (SO 4 ) 3 for ferric sulphate. These might be written Cu!N" 2 O 6 and Fe 2 8 3 O 12 , respectively, but the first formulas are preferable because they enable us to COMBINING PROPORTIONS. 73 see that the compounds are derived /rom nitric acid and sul- phuric acid, HNO 3 and H 2 SO 4 , respectively, while the second formulas do not. The water of crystallization present in many com- pounds is represented by the formula of water (taken the necessary number of times) after the formula of the compound. Thus, while CuSO 4 stands for anhy^ drous cupric sulphate, blue vitriol has the formula Cu- SO 4 . 5 H 2 O ; crystallized zinc sulphate is ZnSO 4 . 7 H 2 O: and crystallized sodium sulphate (Glauber's salt) is Na 2 S0 4 . 10 H 2 0. 75. Combining Proportions. The following list gives the names of some of the more important elements, and the proportions by weight, or a sub-multiple of them, in ivhich these elements unite when they form compounds. s ^ J fc 2 ELEMENT. ELEMENT. i| CO i | N co H Aluminum. Al. 27. Manganese. Mn. 55. Barium. Ba. 137. Mercury. Hg. 200. Bromine. Br. 80. Nitrogen. No 14. Calcium. Ca. 40. Oxygen. 0. 16. Carbon. C. 12. Phosphorus. P. 31. Chlorine. Cl. 35.5 Potassium. K. 39. Copper. Cu. 63. Silicon. Si. 28. Fluorine. Fl. 19. Silver. Ag. 108. Hydrogen. H. 1. Sodium. Na. 23. Iodine. I. 127. Strontium. Sr. 87. Iron. Fe. 56. Sulphur. S. 32. Lead. Pb. 207. Tin. Sn. 118. Magnesium. Mg. 24. Zinc. Zu. 65. A 74 FUNDAMENTAL LAWS. 76. The Use of Combining Proportions. The derivation of the combining proportions given above is of great importance to Chemistry ; however, it depends upon facts which we are not ready to consider at this time. For the present, the combining proportions given in the preceding table will enable us to use equations, formulas, and symbols quantitatively. Thus, if the formula MgO is given to the compound which magnesium forms when it burns in oxygen, a reference to the table will show that in this compound magnesium and oxygen are united in the proportion of 24 grams of magnesium to 16 grams of ox}-gen. From 24 grams of magnesium, therefore, we can get 40 grams of magnesium oxide. Again, if A1 2 O 3 represents correctly the composition of aluminum oxide, the student can see readily that, in this compound, aluminum and oxygen are combined in the proportion of 54 grams of alumi- num to 48 of oxygen. The equation Zn + H 2 SO 4 ZnSO 4 + H 2 thus means, as the result of a simple calculation, that 65 grams of zinc react with 98 grams of sulphuric acid to give 161 grams of zinc sulphate and 2 grams of hydro- gen. 77. How a Compound of Two Elements is Named. A compound of two elements lias the name of each ap- pearing in its name, but the last syllable of the. name of one of the elements is changed to ide. Thus sodium chloride is a compound of sodium and chlorine ; magnesium oxide, one of magnesium and oxygen (here the y HOW TO DISTINGUISH COMPOUNDS. 75 before the ide is omitted) ; calcium carbide, one of calcium and carbon ; and lead sulphide, one of lead and sulphur. As to which element shall have the ending ide, the rule is as follows : If one of the elements is an un- doubted metal, it retains its full name, as in the cases given above, and the ending of the non-metal is changed to ide. If neither is a metal, as in the case of a com- pound of sulphur with oxygen, the name of the one having the smaller resemblance to a metal is changed to ide. Thus, a compound of sulphur with oxygen is called sulphur oxide, but one of sulphur and hydrogen is called hydrogen sulphide. (N. B. Hydrogen has certain metallic properties.) If both elements were metals, the same rule would hold, and the one which possessed the less characteristically metallic properties would have its ending changed to ide. 78. How to Distinguish between Compounds of the Same Two Elements. If there is more than one com- pound of the same two elements, the, ending of the less metallic element is changed in each case, as before, to ide, but the compounds are distinguished from one another in one of two ways. These are, (1) By changing the final syllable of the more metallic (i. e. 9 electro-positive) element ; or, (2) By placing a prefix before the name of the less metallic (i. e.j more electro-uegative) element. 76 FUNDAMENTAL LAWS. By the first method the ending of the metallic ele- ment is changed to ous or ic, as, for example, in the names of the two compounds of mercury and chlorine, which are called mercurowa chloride and mercuric chlor- ide, respectively. The ending ous is given to that one of the compounds which contains the larger proportion of the metallic element (here mercury); the ending ic, on the other hand, to the one containing the smaller proportion of the metallic element. Ous means in chemical phraseology, as in ordinary language, "full of" or "containing much of." The appropriateness of the names for the two compounds of mercury and chlorine will be seen by a comparison of the proportionate amounts of the elements in the two compounds. Thus, mercurous chloride, HgCl, has 200 parts by weight of mercury to 35.5 of chlorine, while mercuric chloride, HgCl 2 , has 200 of mercury to 71 of chlorine. The larger proportion of mercury is, evidently, in the mercurous compound. In the case of some elements, e. g., copper and iron, the Latin names are used, and the ending is applied to these rather than to the English names. Thus, the two common copper oxides are, (a) Cuprous oxide, containing copper, 63 parts, to oxygen, 8 parts ; (6) Cuprtc oxide, containing copper, 63 parts, to oxygen, 16 parts. EXERCISES. 77 Similarly, iron compounds are distinguished by the names, (a) Ferrous oxide, chloride, etc., for the compound contain- ing the larger proportion of iron ; and (6) Feme oxide, chloride, etc., for the one containing the smaller proportion of iron. A second way of distinguishing between two (or more) compounds of the same two elements is to apply a numerical prefix to the less metallic element, and to leave the metallic element unchanged. Thus the names carbon monoxide and carbon o&oxide distinguish the two compounds of carbon and oxygen from each other. The prefix mon means " one " or " first," referring to the formula CO, while di means "two" or "second," referring to the formula CO 2 . In a similar way we distinguish sulphur dioxide, SO 2 , from sulphur trioxide, SO 3 . 79. Exercises. 1. Calculate the percentage composition of sulphuric acid, sodium hydroxide, manganese dioxide (MnO 2 ), and potassium chlorate (KC1O 3 ). 2. What relative quantities of the substances taken and produced are indicated in the equations, 2 KC1O 3 2 KC1 -f 3 O 2 , and 2 K -f- 2 H 2 O > 2 KOH + H 2 ? 78 FUNDAMENTAL LAWS. 3. What quantity of hydrogen could be obtained from the action of 50 grams of potassium upon an excess of water ? 4. Calculate the parts per cent of water of crystallization in blue vitriol, in Glauber's salt, and in gypsum, CaSO 4 . 2 H 2 O. 5. What weight of anhydrous zinc sulphate is contained in 75 grams of the crystallized form, ZnSO 4 . 7 H 2 O ? 6. How many grams of phosphorus must be burned to produce 60 grams of phosphorus pentoxide, P 2 O 5 ? 7. Name the compounds having the formulas CO, CS 2 , ZnCl 2 , Hgl, and BaC 8 . CHAPTER VI. CHLORINE. 80. Existence. Chlorine is a heavy, greenish-yellow gas, of irritating odor and poisonous properties. It was discovered by the Swedish chemist Scheele in 1774, but was not generally considered to be an element until 1809. . Because of its great reactivity, i. e., its tendency to act chemically with other substances, chlorine is not found in nature free, but always in combination with other elements. Its most abundant compounds are sodium, potassium, and magnesium chlorides and hy- drochloric acid, which is hydrogen chloride. Sodium chloride is common salt. 81. Common Method of Preparation. Chlorine is usually prepared by the action of reagent hydro- chloric acid upon manganese dioxide. The apparatus is shown in Fig. 19. A flask containing manganese dioxide (MnO 2 ) in small lumps is provided with a thistle tube and a delivery tube, and is supported so that it may be warmed in a water bath. Con- centrated hydrochloric acid is added through the thistle tube, and the evolution of chlorine begins. The gas is allowed to pass through a wash bottle containing a little water, a drying 79 80 CHLOEINL'. bottle one-third full of concentrated sulphuric acid (a U-tube of calcium chloride may be used instead), and then into a col- lecting bottle, the air of which is to be displaced by chlorine. From the collecting bottle a delivery tube reaches beneath the surface of a solution of sodium hydroxide, which absorbs any H 2 H 2 S0 4 FIG. 19. NaOH escaping chlorine, and enables us to know when the air in the apparatus has been displaced. When a collecting bottle is full of chlorine it is removed and stoppered, and replaced by another bottle until enough gas has been obtained. The bath of hot water is now replaced by one of cold water, and the evolution of chlorine is thus stopped. The proportions by weight of the factors, manganese dioxide and hydrochloric acid, and of the products re- OTUEE METHODS. 81 suiting from their action on one another are as fol- lows : 87 grams manganese dioxide and 146 grams hydrochloric acid (in aqueous solution) give 71 grams chlorine, 126 grams manganous chloride, and 36 grams water. The same facts are represented by the equation, MnO 2 + 4 HC1 = MnCl 2 -f 2 H 2 O + C1 2 . It is probable that the reaction takes place in two stages, and that manganese tetrachloride, MnCl 4 , is formed first and then breaks down into manganous chloride and chlorine. Equations representing these facts are : - - -^^ (1) Mn0 2 -f 4 HC1 = MnCl 4 + 2 H 2 O. (2) The final result, in any case, is that shown in the equation given above. Instead of manganese dioxide and hydrochloric acid, a mixture of common salt, manganese dioxide, and sul- phuric acid is often used. The result is approximately the same, however, for common salt and sulphuric acid give by their action hydrochloric acid. 82. Other Methods. Manganese dioxide is not the only substance that will liberate chlorine from hydro- chloric acid ; potassium dichromate, K 9 Cr 9 O 7 ; potassium chlorate, KC1O 8 ; nitric acid, HNO 3 , and many other substances will do it. 82 CHLORINE. But potassium chlorate and hydrochloric acid give, besides chlorine, an explosive oxide of chlorine. Similarly the chlo- rine formed from the action of nitric acid upon hydrochloric acid is mixed with other substances. Hence these methods are not used to prepare gaseous chlorine. The mixture of nitric and hydrochloric acids is used extensively, however, under the name aqua regia (== royal water) as a solvent for gold, platinum, and other metals not readily attacked by single acids. Aqua regia called, also, nitro-hydrochloric acid is thus only a source of chlorine. Metals dissolved in it are converted into chlorides. It will be noticed that the substances which react with hydrochloric acid to give chlorine are all oxidizing agents. The liberation of chlorine in all the methods described is thus brought about in practically one way, namely, by the oxidation of the hydrogen of hydrochlo- ric acid to water, part of the chlorine being set free. There is a process for the manufacture of crude chlorine on a large scale by the use of atmospheric oxygen as the oxidizing agent ; this is known as Deacon's process. By this method, hydrochloric acid gas mixed with air is passed over heated bricks, which have been soaked in a solution of copper sulphate, or of copper chloride, and then dried. In some way we do not know just how the oxygen is able, under these condi- tions, to act like the oxidizing agents mentioned above. Another method for the production of chlorine and one that may in time displace all others consists in PHYSICAL PROPERTIES. 83 electrolyzing a concentrated solution of hydrochloric acid or of sodium chloride. The chlorine, which is electro- negative, appears at the positive electrode, and the metal or hydrogen at the negative electrode (cf. 37). 83. Physical Properties. Chlorine is about 2% times as heavy as air, and 35.5 times as heavy as hydro- gen. It is easily soluble in water, more than two vol- umes of the gas being absorbed by one of water at the ordinary temperature. The solution called chlorine water possesses many of the properties of the gas. In a warm, saturated solution of common salt, chlorine is only slightly soluble ; it may, therefore, be collected over brine instead of under air. If chlorine is passed into iced water and the solution is cooled below C., a crystalline substance separates out ; this is a compound containing 144 parts of water to 71 of chlorine, and therefore represented by the formula C1 2 . 8 H 2 O. It is called chlorine hydrate. Use may be made of this substance to condense chlorine to the liquid state. For this purpose, the crys- tals of chlorine hydrate are dried between filter papers or on unglazed clay plates, and then put into the closed limb of a tube bent as shown in Fig. 20. The open end is then sealed. The end of the tube containing the chlorine hydrate is now warmed in a 84 CHLORINE. water bath to 30 C., while the other end is surrounded by a freezing mixture of ice and salt. After a short time, liquid chlorine will condense at the drawn-out end. The gas has been liquefied by its own pressure. This experiment can be carried out only in a strong tube securely sealed. Liquid chlorine is an article of commerce ; it is stored and transported in iron cylinders. 84. Chemical Properties. Chlorine is a very dan- gerous substance to inhale, and should, therefore, be generated only in a gas chamber, or where there is a good draught. If it has been taken into the lungs, alcohol or ammonia should be inhaled to counteract, it. It is always well to sprinkle a little ammonia water about in the neighborhood of a chlorine generator. Chlorine is intensely active toward many other ele- ments, forming, by direct union with them, the chlorides. Many substances that combine with oxygen slowly, or not at all at ordinary temperatures, unite readily with chlorine. Powdered antimony and copper foil (the lat- ter must be hot) glow when put into chlorine, the prod- ucts being antimony trichloride (SbCl 3 ) and cupric chloride (CuCl 2 ) respectively. Sodium, tin, magne- sium, and phosphorus all give corresponding chlorides when put into the gas. But it is toward hydrogen that chlorine shows its most remarkable behavior, for while the two gases do not combine at all in the dark, and only very slowly in diffused light, yet they unite with explosive violence in sunlight. ACTION OF CHLORINE AND AMMONIA. 85 The mixture of hydrogen and chlorine may also be exploded by a burning match or by the electric spark. Chlorine shows this tendency to combine with hydro- gen not only when the hydrogen is in the free state, but also when it is united with other elements. As illus- trations we may take the action of chlorine toward water, ammonia, and turpentine. 85. Action of Chlorine and Water. The aqueous solution of chlorine may be preserved for a long time if kept cold and in the dark, but it decomposes rapidly in sun- light, giving as final products hydrochloric acid and oxy- gen. The equation is : 2H 2 O + 2 C1 2 4HC1 + O 2 . (See, however, " bleaching powder.") If the decomposition of chlorine water by sunlight is carried out in a long tube (Fig. 21), a colorless gas will collect in the upper part of the tube. The gas is oxygen. Much , of this gas will also be found in the solution. Oxygen Chlorine Water Chlorine Water FIG. 21. 86. Action of Chlorine and Ammonia. When the hydrogen of ammonia is appropriated by chlorine, hydro- chloric acid and nitrogen are formed, as represented by the equation, 2 NH 3 + 3 C1 2 6 HC1 + N a . With the excess of ammonia the hydrochloric acid gives 86 CHLORINE. ammonium chloride, NH 4 C1, the material of the white smoke seen when ammonia and chlorine gases come together. 87. Action of Chlorine and Turpentine. The behav- ior of chlorine and turpentine may be shown by immers- ing a piece of filter paper in warm turpentine and then plunging it into a jar of chlorine. The turpentine soon ignites, and burns with a smoky flame. Turpentine is a compound of carbon and hydrogen, and the chlorine, by uniting with some of the hydrogen, but not with the car- bon, sets the carbon free in the form of a dense, black smoke. 88. Uses of Chlorine. Chlorine is used in large quantities as a bleaching and disinfecting agent, and is generally made for these purposes from bleaching powder, or " chloride of lime." Bleaching powder is a white substance formed by the action of chlorine upon "slaked" lime (calcium hydroxide), and is easily de- composed by acids, even by the carbon dioxide of the air, with evolution of chlorine. Fabrics to be bleached by the chlorine process are, therefore, immersed in a bath of dilute acid, and then in one of chloride of lime. In this way chlorine is set free in immediate con- tact with the coloring matter of the cloth, and bleaches it. Chlorine is not a bleaching agent ordinarily, unless water is present, hence it is likely that chlorine itself does not act upon the coloring matter, but upon the water. As a result, oxygen is probably set free ; and it is oxygen, and not chlorine, that bleaches the cloth. Since ordinary oxygen is not able to effect this change, we assume that oxygen at the instant of its libera- tion from water is in a condition different from that in which EXERCISES. 87 we ordinarily find it. We say that it is in a nascent condition. The reason for the peculiar behavior of an element in its nas- cent state will be given later. Compared with the old bleaching process, which con- sisted in exposing the fabric to the oxidizing agents of the air, the chlorine method is of course very rapid, but, unfortunately, the bleaching agent used too often attacks the fiber of the cloth as well as its coloring matter. Hence delicate materials, such as the better grades of straws, laces, silks, and woolens, are usually decolorized by sulphur dioxide, which, although it does not bleach so permanently as chlorine, has yet the ad- vantage of acting less upon the fabric. The action of chloride of lime as a disinfectant is similar to its action as a bleaching agent : nascent oxy- gen is formed, and this destroys the micro-organisms of the surrounding air. 89. Exercises. 1. How many grams of chlorine can, theoretically, be ob- tained by the electrolysis of 50 grams of hydrochloric acid ? 2. How many grams of manganese dioxide are required to give with an excess of hydrochloric acid 10 grams chlorine ? 3. What will be the volume of 40 grams chlorine Under con- ditions at which 1 liter of hydrogen weighs 0.09 gram ? 4. How much silver chloride, AgCl, can be formed by burning 54 grams silver in chlorine gas ? 5. Calculate the per cent of chlorine in sodium chloride. 6. How many liters of chlorine can be made by the action of 25 grams of manganese dioxide with an excess of hydro- chloric acid ? (Assume that 1 liter of chlorine weighs 3 grams.) CHAPTER VII. HYDROCHLORIC ACID. 90. Existence. Hydrochloric acid is a colorless, heavy gas which fumes in moist air and dissolves readily in water. It is found in only small amounts in nature, e. g., in volcanic gases and in some springs. It makes up about 0.02 of 1^ of the gastric juice. 91. Preparation: Common Laboratory Method. Hydrochloric acid may be prepared readily by the action of sulphuric acid, H 2 SO 4 , upon common salt, NaCl. The apparatus is shown in Fig. 22. FIG. 22. The flask contains common salt and sulphuric acid diluted with half its volume of water, and cooled. (Caution! In di- COMMERCIAL MANUFACTURE OF THE ACID. 89 luting sulphuric acid we always pour the acid into the water.) A hot-water bath serves to heat the flask. The first bottle contains concentrated sulphuric acid to dry the gas, and the second bottle is the collecting vessel. A beaker of water collects any escaping gas and shows when the gas in the col- lecting bottle is free from air. The reaction between common salt and sulphuric acid takes place according to the equation, NaCl -f H 2 SO 4 > ^TaIISO 4 -f IIC1 ; that is to say, 58.5 grams of sodium chloride and 98 grams of sulphuric acid give 120 grams sodium hydro- gen sulphate and 36.5 grams hydrochloric acid. If there is an excess of sulphuric acid, the sodium chloride is all used up, and the white solid which crystallizes in the generating flask when the latter cools is sodium hydrogen sulphate. 92. Commercial Manufacture of the Acid. In the first stage of the manufacture of sodium carbonate (soda) by the " Le Blanc " process, sodium chloride is treated with sulphuric acid, with the results illustrated by the equation just given. Since, however, sodium sulphate (Na 2 SO 4 ) and not sodium hydrogen sulphate is the product wanted, the sodium hydrogen sulphate is converted into sodium sulphate by heating it with more sodium chloride to a high temperature. The reaction which then takes place is represented thus : INaCl -f NaHSO, Na 2 SO 4 -f- HC1. 58.5 g. 120 g. 142 g. 36.5 g. 90 HYDROCHLORIC ACID. From this we see that the hydrogen of sulphuric acid, like that of water, may be replaced in two stages ; for while at low temperatures only 58.5 grams of sodium chloride react with 98 grams of sulphuric acid, at a high temperature a second quantity of sodium chloride, equal to the first, is able to react, and thus produces a second quantity of hydrochloric acid. This will be more evi- dent when the two equations are written together : (1) NaCl -f g > S0 4 - g a > S0 4 + HC1. (2) NaCl + a > S0 4 - > S0 4 + HC1. The hydrochloric acid formed as a by-product in the soda manufacture is conducted into water, and the solu- tion is sold as commercial hydrochloric acid. It is usu- ally somewhat colored by slight impurities. 93. Physical Properties. Hydrochloric acid is about li times as heavy as air, and 18^ times as heavy as hydrogen. It is very soluble, 505 c.c. being held by 1 c.c. water at C. and 760 mm. pressure. The con- centrated solution of the pure gas in distilled water at the ordinary temperature is the " chemically pure " (c. p.) reagent hydrochloric acid. This is a colorless liquid of specific gravity 1.2 (Water = 1). The concentrated solution of hydrochloric a,cid fumes strongly in moist air because the escaping gas condenses some of the water vapor of the air. COMPOSITION OF HYDROCHLORIC ACID. 91 The dry gas can be converted into a colorless liquid at a low temperature and great pressure. 94. Volumetric Composition of Hydrochloric Acid. - The composition of hydrochloric acid may be demon- strated in the same way as that of water, viz., by electrol- ysis. If an electric cur- rent of sufficient strength is passed through a concen- trated aqueous solution of the acid, hydrogen is pro- duced at the electrode, and chlorine at the -f- elec- trode. Cl Zn FIG. 23. The gases collect at un- equal rates, at first, because of the greater solubility of the chlorine ; but when the liquid has become saturated with both gases the hydrogen and the chlorine gather in the collecting tubes at the same rate. See Fig. 23. Therefore hydrochloric acid gas must be composed of hydrogen and chlorine united in equal proportions by volume. The same fact may be proved synthetically, for if a mixture of hydrogen and chlorine be exploded, it will be found that equal volumes of the two gases have dis- appeared. The volume of hydrochloric acid formed will be equal to the sum of the uniting gases. These facts may be represented graphically as follows : 92 HYDROCHLORIC ACID. n + n > nn 1 vol. hydrogen. 1 vol. chlorine. 2 vols. hydrochloric acid. Note that this case is different from that of water ; for, in the production of 2 volumes of steam, 2 volumes of hydrogen united with 1 of oxygen, i. e., 3 volumes of the mixed gases gave only 2 volumes of the product. Since the weights of equal volumes of hydrogen and chlorine are about as 1 : 35.5, the two volumes of hydro- chloric acid formed by their union should be about 36.5 times as heavy as one volume of hydrogen. Hence hy- drochloric acid should be about 18.25 (= 36.5 2) times as heavy as hydrogen. This is actually the case. That hydrochloric acid gas gives, when decomposed, one-half of its own volume of hydrogen, may be shown by the action of sodium. For con- venience the sodium is diluted by alloying it with mercury. A small amount of the resulting sodium amalgam is put into a long measuring tube (Fig. 24) full of hydrochloric acid gas. The open end of the tube is then closed with the thumb, and the tube is shaken /Sodium Amalgam vigorously. When the thumb is removed thumb- under water, some of the water rushes FIG. 24. U P m ^ the tube to replace the chlorine. If the water inside and outside the tube is now brought to the same level, the volume of the residual hydrogen will be found to be approximately one-half that oi the hydrochloric acid taken. CHLORIDES. 93 95. Acid Properties. In addition to the properties already given, hydrochloric acid (i. e., the aqueous solu- tion or the moist gas) has the characteristics of a special class of substances called acids. Acids may be de- scribed roughly as having a sour (acid) taste, the ability to change certain vegetable colors, e. g., blue or purple litmus to red, and also the power to neutralize the prop- erties of another class of substances, viz., the bases. Thus the action of hydrochloric acid upon sodium hy- droxide (a base) gives sodium chloride (common salt) and water. Metals, too, react with acids, forming salts. Thus, zinc with hydrochloric acid gives zinc chloride (a salt) and hydrogen. These reactions are shown in the equations, NaOH + HC1 - N"aCl -f H 2 O, and 40 36.5 58.5 18 Zn + 2 HC1 - ZnCl 2 + H 2 . 65 73 136 2 The properties of acids, bases, and salts will be considered in the next chapter. Hydrochloric acid is one of the most important acids. It is made on a large scale, and is used in enormous quantities. 96. Chlorides. The chlorides may all be considered hydrochloric acid with its hydrogen replaced by a metal ; the acid itself is often called hydrogen chloride. The most important chlorides have been given in 80 ; 94 HYDROCHLORIC ACID. others are barium chloride, BaCl 2 , silver chloride, AgCl, and ferric chloride, FeCl 3 . The most abundant chlo- ride is, of course, common salt, NaCl. The chlorides of most of the common metals are soluble in water. Exceptions are silver chloride, AgCl, and mercurous chloride, HgCl. Lead chloride, PbCl 2 , is only slightly solu- ble in cold water, but more readily in hot water. When, therefore, solutions of salts of silver, lead, and mercury (in its mercuroMS condition) are treated with a solution of a chloride, the chlorides of these metals are precipitated. 97. Exercises. 1. 300 c.c. of hydrogen and 250 c.c. of chlorine were mixed and exploded. What was the product ? Its volume ? Which of the gases used was in excess ? How much ? 2. Calculate the percentage composition of hydrochloric acid? 3. How many grams of sodium chloride are needed to yield, with sulphuric acid, 20 grams hydrochloric acid gas ? 4. How many grams hydrochloric acid can be made from 35 grams potassium chloride, KC1 ? 5. What weight of sodium chloride is necessary to produce, with sulphuric acid, 20 liters of hydrochloric acid gas when 1 liter of hydrochloric acid gas weighs 1.63 grams ? CHAPTER VIII. ACIDS, BASES, AND SALTS. 98. Acids. In all of our study of Chemistry we shall have to deal constantly with bodies belonging to the classes adds, bases, or salts. Let us first consider some of 'the acids. One of these, viz., hydrochloric acid, we have already studied at some length; other important acids are nitric acid, sulphuric acid, acetic acid, arid tartaric acid. Only a short description of these will be given here. Nitric acid, HNO 3 , is a colorless liquid. It ordinarily has a sharp odor and is very corrosive in concentrated form. It turns the skin yellow. A dilute solution of nitric acid is sour, turns blue litmus and neutral litmus pink, decomposes carbonates, e. g., marble, or calcium carbonate, and acts upon many metals. Sulphuric acid, H 2 SO 4 , is a heavy, oily liquid which dis- solves in water with the evolution of much heat. It chars organic substances, and therefore becomes dark colored when exposed for a time to the dust of the ah-. Its dilute solution has a sour taste, and acts upon litmus and carbon- ates as nitric and hydrochloric acids do. With many metals dilute sulphuric acid gives sulphates and hydrogen. Acetic acid, HC 2 H 3 O 2 , is a colorless, sharp- smelling liquid. Like the other acids, it has a sour taste, acts upon 96 ACIDS, BASES, AXD SALTS. litmus, and decomposes carbonates. Vinegar is a dilute solution of acetic acid. Tartaric acid, H 2 C 4 H 4 O 6 , is a white, crystalline solid, soluble in water. Its solution has properties similar to those of the other acids. All of the acids named, except hydrochloric acid, may be looked upon as water joined to an oxide. The oxide is called the anhydride of the acid. Thus nitrogen pentoxide, N 2 O 5 , is the anhydride of nitric acid, for N 2 O 5 -f H 2 O = 2 HNO 3 . Similarly, sulphur trioxide, SO 3 , is the anhydride of sulphuric acid. The most important property of all acids is their power of reacting with the hydroxides of metals. When a solution of any of the above acids is treated with a solution of a metal hydroxide, e. g., sodium hy- droxide, an evolution of heat takes place, and if the correct amount of sodium hydroxide is used, the taste of the acid, its power to change litmus, and to act upon carbonates and metals, will all disappear. The acid has been neutralized by the sodium hydroxide. 99. Bases. The general properties of sodium hy- droxide and its relation to water have already been, given (y any pressure, however great. Indeed, pressures up to 3,600 atmospheres have been applied without avail. This is due to the fact that there is for every gaseous substance _ 116 NITEOGEN AXD THE ATMOSPHERE. a maximum temperature above which the gas cannot be liquefied; this is called the critical temperature of the gas. A "permanent" gas thus differs from an easily condensible gas in this respect, viz., that the critical temperature of a condensible gas is above the ordinary temperature, while that of the permanent gas lies far below the ordinary temperature. Such gases as air, hydrogen, etc., are, therefore, condensed only at a very low temperature and great pressure. Two general methods are used to liquefy true gases. In the first method the gas to be condensed is first cooled to its critical temperature, and is then subjected to pressure. In the second method the gas is first strongly compressed, and is then cooled to its critical tempera- ture. The second method has been used recently to liquefy air on a commercial scale. The apparatus used consists, essentially, of two systems of pipes, the pipes of the outer system forming a jacket surround- ing those of the inner system. The air of the inner system is that to be liquefied. By means of compressing engines (c/. Fig. 34), air is forced into each system under great pressure. Much heat is thereby evolved. When the compressed air has cooled to the ordinary temperature, some of the air of the outer pipes is allowed to escape. The sudden expansion of the air remaining in the outer pipes then causes an amount of heat to be absorbed which is equal to that evolved when the compression took place. The air of the outer pipes, now intensely cold, cools the compressed air of the inner pipes below the critical temperature of air, and thus liquefies it. THE PROPORTION OF OXYGEN IN AIR. 117 122. Properties of Liquid Air. In the liquid con- dition air is colorless, has about the same density as water, and boils at 190 C., under ordinary pressure. When freshly made, liquid air is about half oxygen, but the proportion of oxygen increases by the evaporation of the nitrogen (liquid nitrogen boils about 10 C. below liquid oxygen) until the liquid is over 90^? oxygen. Liquid air is preserved in open, double-walled vessels called Dewar bulbs (Fig. 29). The space between the walls of the bulbs is exhausted of air to secure non-conductivity of heat. Tin or CN /O wooden boxes having double walls filled with silk may also be used. Alcohol, liquid carbon dioxide, mercury, etc., solidify when placed in liquid air, and steel burns in it like tinder ; yet the hand may be held in it for a short time with- out injury, because protected by a non-conducting film of air in the gaseous state. 123. Determination of the Proportion of Oxygen in Air. The amount of oxygen in a given volume of air is usually determined either (1) by absorbing the oxygen, or (2) by exploding the air with a known vol- ume of hydrogen. The phosphorus absorption method, a crude form of 118 NITROGEN AND THE ATMOSPHERE. which was described under nitrogen ( 4 H 2 O + 2 NO. Copper and nitric acid of specific gravity 1.2 do not react according to the equation, Cu -f- 2 HN0 3 = Cu(NO,) 2 + H 2 , for the reaction is an oxidation : 3 Cu + 2 HN0 3 3 CuO + 2 NO + H 2 O. Copper nitrate is indeed formed, 63 grams of copper giving 124 grams of copper nitrate, but no hydrogen is evolved. Instead of hydrogen we get a gas which, though colorless itself, forms brown fumes when it comes in contact with oxygen. This gas is nitric oxide, NO, and the brown fumes consist of nitrogen dioxide, NO 2 ; hence the copper must have reduced some of the nitric acid. The cupric oxide then reacts with the nitric acid, giving cupric nitrate and water. 3 CuO + 6 HNO 3 3 Cu(NO 3 ) 2 + 3 H 2 O. 154 NITROGEN ACIDS AND OXIDES. The complete equation is, therefore, 3 Cu + 8 HNO 3 = 3 Cu(NO 3 ) 2 + 4 H 2 O + 2 NO. The equation for the action of silver and nitric acid is simi- lar, viz. , 3 Ag + 4 HNO 3 = 3 AgNO 3 -f 2 H 2 O + NO. 160. Aqua Regia. Gold and platinum alone of all the more common metals do not react with nitric acid. As was stated under Chlorine (cf. 82), these two metals are soluble in aqua regia and in chlorine water. Aqua regia is made by mixing nitric acid with hydro- chloric acid (three volumes of hydrochloric acid to one of nitric acid) ; it is merely a source of nascent chlorine. Aqua regia is used to dissolve many other metals be- sides gold and platinum. 161. Oxidation by Nitric Acid. It is not only to- ward metals that nitric acid acts as an oxidizing agent. Phosphorus is converted by it into phosphoric acid, and sulphur into sulphuric acid ; while glowing charcoal burns in the acid much as it would in oxygen itself. Organic coloring matters, e. g., indigo solution, are oxi- dized by nitric acid to colorless bodies. The oxidizing power of nitric acid is much greater if some of the higher oxides of nitrogen are present. Such an acid is fuming nitric acid, a red liquid containing much nitrogen di- oxide, NO 2 . This acid is used when very rapid oxidation is desired. 162. Explanation of Oxidation by Nitric Acid. - The student will be able to understand the peculiar FOEMATION OF NITRATES IN NATURE. 165 behavior of nitric acid better if he will consider nitric acid as made up of water and nitrogen pentoxide , N 9 O., as represented by the equation, He will then see that the oxidizing action of nitric acid is due to the loss of oxygen by nitrogen pentoxide. Even without a reducing agent, nitrogen pentoxide readily gives nitrogen tetroxide and oxygen (cf. 158) ; but in the presence of such an agent the decomposition of nitrogen pentoxide is very easy indeed. Nitrogen tet- roxide, N 2 O 4 ; nitrogen trioxide, N 2 O 3 ; nitric oxide, NO ; nitrous oxide, N 2 O ; nitrogen, N 2 , or even ammo- nia, NH 3 , may be formed, according to circumstances. The following equations represent some of the reactions that may take place : N 2 O 5 = N 2 O 4 -f- O. This takes place when concentrated nitric acid is heated, or treated with metals. N 2 O 5 = N 2 O 3 -j- 2 O. This takes place with starch, arsenic trioxide, etc., and acid of specific grav- ity 1.3. N 2 O 5 = 2 NO -f 3 O. This takes place with copper, silver, etc., and acid of specific gravity 1.2. NgOg = N 2 O -|- 4 O. This takes place with zinc and acid of specific gravity 1.1. 163. Formation of Nitrates in Nature. We have already learned (c/1 112) that nitrogen, oxygen, and water vapor are combined by the action of the electric spark into nitric acid ; consequently, it is very probable that in some regions much nitric acid is formed in this * 156 NITROGEN ACIDS AND OXIDES. way and gets into the soil. All the nitric acid of com- merce, however, is made from the nitrates; these are found as natural deposits in certain places. The " nitrate beds " were probably formed by the oxidation, through the agency of bacilli, of nitrogenous organic matter in the presence of alkali. The most ex- tensive deposits of nitrates in the world are those of the Atacama Desert in Chile. The alkali present is sodium carbonate, hence Chile saltpeter is sodium nitrate and not the potassium salt. 164. Manufacture of Potassium Nitrate. Potas- sium nitrate was formerly obtained almost exclusively from Asiatic countries, where it appeared as a deposit on the ground; but nowadays most of it is made from sodium nitrate. The process may be carried out as follows : Hot, fairly concentrated solutions of potassium chloride and sodium nitrate are mixed, and the resulting solution is poured off from the crystalline deposit of sodium chloride which sepa- rates out. The solution contains much potassium nitrate and small amounts of sodium chloride, etc. This impure potassium nitrate is then redissolved in as small an amount of hot water as possible, and allowed to crystallize out ; by several recrystal- lizations pure potassium nitrate is obtained. The equation representing the formation of potassium ni- trate is, KC1 -f NaKOg > KKO 3 + KaCl. It applies only to a mixture of concentrated solutions of the factors. USES OF NITRIC ACID AND THE NITRATES. 157 In certain European countries the farmers cultivate niter by introducing the proper micro-organisms into a mixture of alkali and nitrogenous matter. 165. Uses of Nitric Acid and the Nitrates. - Nitric acid has many important applications : nitro- benzene, glyceryl nitrate (nitroglycerine), and nitrates of cellulose (collodion, gun cotton, celluloid, etc.) are made by its agency. The nitrates are used in the manufacture of gun- powder and various explosives, and in the preservation of meats. Glyceryl nitrate (wrongly called nitroglycerine) is made by the action of a mixture of concentrated nitric and sulphuric acids at a low temperature upon glycerine. It is a thick, greenish oil of a very unstable nature, and very explosive. It is usually mixed with a porous earth, and appears in the mar- ket chiefly as dynamite. Gun cotton and collodion are made by the action of a mixture of nitric and sulphuric acids upon cotton. Celluloid is a mixture of gun cotton and camphor. Nitrobenzene is made from benzene, C 6 H 6 , and the nitric- sulphuric acid mixture. When nitrobenzene is reduced by nascent hydrogen, it gives aniline, the starting material in the manufacture of the aniline dyes. Potassium nitrate is important chiefly as a constituent of gunpowder, which, as was stated in 26, is a physi- cal mixture of potassium nitrate with sulphur and charcoal in various proportions. The exact amount of each of the constituents depends upon the purpose for 158 NITROGEN ACIDS AND OXIDES. which the powder is to be used, but the function of the potassium nitrate is always the same, viz., to furnish oxygen for the combustion of the sulphur and the charcoal. When gunpowder is ignited, it forms gases that occupy at ordinary pressure several hundred times the volume of the original gunpowder. The reactions taking place in the explosion of one kind of gunpowder are probably represented by the equation, 2 KNO 3 + C + S K 2 SO 4 -f CO 2 -f N 2 . The chief use of potassium nitrate as a preservative is in the preparation of " corned " beef. 166. Nitrogen Pentoxide. Nitrogen pentoxide (" penta " = five) is of theoretical importance because it is the anhydride of nitric acid, i. e., it is nitric acid minus water. The relation of nitrogen pentoxide to nitric acid is evident from the equation, N 2 5 + H 2 2 HN0 3 . Nitrogen pentoxide is a white, crystalline solid, made by the distillation of anhydrous nitric acid with phosphorus pen- toxide. 167. Nitrous Acid. Nitrous acid is probably con- tained in a solution of nitrogen trioxide in water ; ni- trogen trioxide is, therefore, called nitrous anhydride, just as nitrogen pentoxide is nitric anhydride. This re- lation is shown by the equation, N 2 3 + H 2 2 HN0 2 . NITROGEN TE I OXIDE. 159 The meaning of the ending ons as applied to nitrous acid has already been given (cf. 106). Nitrous acid itself has not been made ; but its salts (called nitrites to distinguish them from the salts of nitric acid) are formed when a solution of nitrogen trioxide is neutralized by bases. Thus, potassium hydroxide and a solution of nitrogen tri- oxide give potassium nitrite and water, as is shown by the equation, 2 KOH + 1SLO, - 2 KXCX 4- II O. I J o ti I 2 If we consider the solution of nitrogen trioxide to be nitrous acid, the equation becomes, KOII -f Iim) 2 = K^s T O 2 -f II 2 O. A second way in which nitrites are formed is by the abstraction of oxygen from nitrates. Thus, potassium nitrate loses one-third of its oxygen when heated to a high temperature. The equation is, The temperature required is much lower if the nitrate is mixed with lead when heated ; for the lead unites with the liberated oxygen to form lead oxide, PbO. 168. Nitrogen Trioxide (N 2 3 ). As was stated in 167, nitrogen trioxide is the anhydride of nitrous acid. It may be made by the action of nitric acid of specific gravity 1.3 upon arsenic trioxide or upon starch. The brown fumes called the trioxide may be condensed by a freezing mixture to a blue liquid, which is the real trioxide. When, however, this liquid is distilled, the vapors are not ni- 160 NITROGEN ACIDS AND OXIDES. trogen trioxide, but a mixture of the dioxide (NO 2 ) with nitric oxide. *i69. Nitrogen Dioxide and Nitrogen Tetroxide (N0 2 and N 2 4 ). Nitrogen dioxide lies between nitro- gen trioxide and nitrogen pentoxide as regards the pro- portion of oxygen it contains. Below 22 C. it is a liquid ; but it is usually known in the form of its vapor. Nitrogen tetroxide (" tetra " four) exists at low temperatures, but it dissociates readily, to some extent even at C., into nitrogen dioxide. This is shown in the equation, N 2 4 = 2 N0 2 . The degree of the dissociation is shown by the dark- ening of the color, nitrogen tetroxide being colorless, but nitrogen dioxide brown. When nitrogen tetroxide is dissolved in much cold water, the solution contains a mixture of nitrous arid nitric acids ; hence nitrogen tetroxide is the anhydride of both of these acids. The equation showing this fact is, N 2 O 4 (= 2 NO 2 ) -f H 2 O HNO 2 + HXO 3 . Nitrogen tetroxide is formed when lead nitrate, Pb(NO 3 ) 2 , is heated. The vapors may be condensed by passing them through a U-tube surrounded by a freezing mixture. The decomposition of lead nitrate corresponds exactly with that of nitric acid, as is shown by the equation, Pb(N0 3 ) 2 = PbO + N 2 4 + O (_of. 158). *The names used are in accordance with the latest nomenclature. NITRIC OXIDE. * 161 170. Nitric Oxide (NO). Two other oxides of ni- trogen are known, viz., nitric oxide and nitrous oxide ; .both are colorless gases at the ordinary temperature and pressure. Nitric oxide is produced when nitric acid of specific gravity 1.2 is allowed to react with copper. The equa- tion has already been given (N 2 + H./>. Nitrous oxide may thus be looked upon as in some sense the anhydride of hyponitrous acid ; but the union of nitrous oxide and water to form the acid does not take place. 173. Exercises. 1. How many grams of nitric acid can be made, theoreti- cally, from 1 kg. sodium nitrate with sulphuric acid ? 2. What will be the volume of 28.4 grams of commercial nitric acid ? 164 NITROGEN ACIDS AND OXIDES. 3. How much sulphuric acid (calculated as 100% H 2 8O 4 ) will be needed to give 1,260 grams of nitric acid with potassium nitrate, if potassium hydrogen sulphate is formed ? 4. Calculate the percentage composition of nitric acid ? 5. How much potassium hydroxide is required to neutralize exactly a solution containing 42 grams nitric acid ? 6. How could you separate a mixture of silver and gold chemically ? 7. How many grams nitrogen tetroxide could be made from 450 grams lead nitrate ? How much oxygen ? 8. How many grams nitric oxide can be made, theoretically, from 100 grams nitric acid with copper? How much copper is needed? How much cupric nitrate is formed? 9. How many liters of nitrous oxide at C. and 760 mm. can be made from 240 grams ammonium nitrate ? How many at 20 C. and 720 mm. ? 10. How would you distinguish between nitrous oxide and oxygen ? n. How many cubic centimeters of each of its constituents combine to form 100 c.c. nitric oxide ? To form 100 c.c. of nitrous oxide ? 12. Name three different classes of nitrates, basing the dif- ference upon the way in which the members of each class decompose when heated. CHAPTER XIII. SULPHUR AND ITS COMPOUNDS. 174. Occurrence and Preparation of Sulphur. SuL phur occurs in nature in both a free and a combined form. In the free condition it is obtained chiefly from Sicily, Mexico, and, to some extent, from Louisiana. Natural sulphur is usually found mixed with much earthy material, from which it must be separated to pre- pare it for the market. The first operation in the purification of sulphur usu- ally consists in heating the natural product ; the sul- phur melts and flows away, leaving the infusible im- purities behind. In the second operation the partially purified sul- phur is distilled from large iron retorts (see Fig. 38), and is thus separated from FIG . gg. less volatile impurities. The melted sulphur in the reservoir A is allowed to flow from time to time into the retort 5, in which the sulphur is vaporized. The sulphur vapor which passes into the con- 165 166 SULPHUR AND ITS COMPOUNDS. denser collects either in the liquid state, at the bottom ((7) of the condenser, or in a solid state upon the cold walls (D). The liquid sulphur is run into molds to crystallize, thus producing the " roll-sulphur," or " brimstone " of commerce ; the sul- phur which solidifies upon the walls appears in the form of fine meal and is called "flowers " (more correctly, "flour") of sulphur. 175. Physical Properties. Sulphur, like many other elements, exists in several different physical forms; consequently, in giving the properties of sul- phur we must specify the kind of sulphur to which we are referring. The several varie- ties of sulphur may he grouped into three classes : (1) Ordinary, or rhombic, sul- phur (Fig. 39). This is the form that occurs in nature. All other forms revert to this form. Its spe- FIG ' 39> cine gravity is 2.07. (2) Prismatic sulphur (Fig. 40). This is formed by the slow cooling of fused sulphur of any of the other varie- ties. Its specific gravity is 1.96. (3) Amorphous sulphur. This form is produced when sulphur at tempera- tures above 230 C. is chilled rapidly, as by pouring it into cold water. At ordinary temperatures amorphous sulphur changes slowly (after some days) into the ordinary form ; at about CHEMICAL PROPERTIES. 167 100 C., however, the change is instantaneous, and much heat is evolved. Ordinary and prismatic sulphur are readily soluble in carbon disulphide, CS 2 ; but amorphous sulphur is partly in- soluble in that liquid. The form known as " flowers" of sulphur is both crystalline and amorphous. It consists of a crystalline kernel and an amorphous covering. The existence of an element in several forms is called allotropism, and the different varieties of the element are called its allotropic forms ( FeSO 4 -f H 2 S, and FeS + 2 HC1 FeCl 2 -f H 2 S. The iron sulphate or chloride formed remains in solu- tion. The method just described gives hydrogen sulphide in a form good enough for ordinary use, but not pure. Pure hydro- gen sulphide is prepared by the action of concentrated hydro- chloric acid upon antimony trisulphide, Sb 2 S 3 . The equation is, Sb 2 S 3 -f 6 HC1 2 SbCl 8 + 3 H 2 S. 1 80. Properties of Hydrogen Sulphide. Hydrogen sulphide has the odor of rotten eggs. It is formed by the decomposition of most organic substances containing sulphur. So-called "sulphur" waters OAve their proper- ties to the hydrogen sulphide dissolved in them. The gas is 1.18 times as heavy as air. Hydrogen sulphide is very soluble in water; one volume of water absorbs at standard conditions three volumes of the gas. The aqueous solution is readily decomposed, especially in the light and when warm, by the oxygen of the air. The hydrogen is thus converted SULPHIDES. 171 into water, and the sulphur is set free ; consequently a solution of hydrogen sulphide soon loses its odor and deposits sulphur. The equation is, 2 H 2 S + (X - > 2 HoO + 2 S. Similar to the action of oxygen is that of chlorine, which forms with hydrogen sulphide or its aqueous solution hydro- chloric acid and sulphur, according to the equation, II 2 $ + Cl, - 2 HC1 + S. Iodine acts in the same way, viz.: This method is used for the preparation of hydriodic acid. Hydrogen sulphide is, as we might expect, a reducing agent. 181. Sulphides. Hydrogen sulphide is a weak acid, and is therefore called hydro sulphuric acid. Its salts, the sulphides, may be formed in several ways ; these are, (1) By the reduction of a sulphate or sulphite ; (2) By the neutralization of an alkali with hydrogen sulphide ; (3) By the addition of hydrogen sulphide to a soluble salt of the metal whose sulphide is to be formed. An illustration of the first method is the reduction of sodium sulphate, when heated with charcoal, to sodium sulphide. The equation is, Na 2 S0 4 + 4 C - > Na 2 S + 4 CO. An illustration of the second method is the absorption 172 -SULPHUB AND ITS COMPOUNDS. of hydrogen sulphide by a solution of sodium hydroxide. The equation is, 2 KaOH -f H 2 S > Xa 2 S + 2 H 2 O. Hydrosulphides. The equation just given represents only the final products formed in the ordinary method of 'preparing sodium sulphide ; for, if we pass hydrogen sulphide into sodium hydroxide solution until no more hydrogen sulphide is absorbed, we obtain sodium hydrosulphide,N&SH. The equation is, KaOH -f H 2 S = NaSH + H 2 O. If we now add to the sodium hydrosulphide as much sodium hydroxide as we used originally, we shall obtain sodium sul- phide, Na 2 S, as is represented in the equation, NaOH -f NaSH = !Na 2 S + H 2 O. To prepare ammonium sulphide, N"II 4 ) 2 S, we employ a simi- lar method. 182. Precipitation of Sulphides. The third method of forming a sulphide, viz., by adding hydrogen sul- phide to a soluble salt of the metal, will succeed only in case the sulphide desired is insoluble in the solvent present. Thus if hydrogen sulphide, either in gaseous form or in aqueous solution, is added to a solution of cupric sul- phate, CuSO 4 , cupric sulphide and sulphuric acid are formed, according to the equation, CuSO 4 + H 2 S > CuS -f H 2 SO 4 . The cupric sulphide will appear as a black precipitate. In future we shall usually italicize the formula of a precipitate, PRECIPITATION OF SULPHIDES. 173 Explanation. The student will notice that in every case the action of hydrogen sulphide upon the salt of a metal must tend to produce a free acid in addition to the sulphide of the metal. Now, we know that free acids generally act upon sulphides giving hydrogen sulphide, this being in fact the way in which hydrogen sulphide is prepared (cf. 179) ; hence the precipitation of cupric sulphide from a solution of cupric sulphate by hydrogen sulphide must be possible only because the cupric sulphide is insoluble in ' the dilute acid formed at the same time. Sulphides that are soluble in dilute acids cannot, therefore, be precipitated, or only incompletely, by hydrogen sulphide. Thus, no precipitate is produced when hydrogen sulphide is added to manganese sulphate solution, because the reverse re- action, represented by the equation, MnS + H 2 SO 4 MnSO 4 -f H 2 S, is the one that tends to take place. If, however, we use, in- stead of hydrogen sulphide, a soluble salt of hydrogen sulphide, precipitation of manganese sulphide occurs, for the reaction cannot produce free acid. Thus, with sodium sulphide the reaction is represented by the equation, MnSO 4 -f Na 2 S > MnS -f Na 2 SO 4 . Some sulphides, however, are soluble in water itself ; such sulphides cannot, of course, be precipitated by either hydrogen sulphide or its salts. Barium and calcium sulphides are ex- amples. 174 SULPHUR AND ITS COMPOUNDS. 183. Carbon Bisulphide (CS 2 ). Carbon disulphide is formed by the direct union of carbon and sulphur, the usual method being to pass sulphur vapor over hot charcoal. When pure, carbon disulphide is colorless and has an ethereal odor ; but as obtained commercially it is often yellow and has a disagreeable smell. The liquid boils at 47 C. Carbon disulphide is very inflammable. The prod- ucts of its combustion are carbon dioxide and sulphur dioxide, as represented by the equation, CS 2 +30 2 = C0 2 + 2S0 2 . The chief use of carbon disulphide is as a solvent for sul- phur, caoutchouc, phosphorus, iodine, etc. 184. Manufacture of Sulphuric Acid. - - The mod- ern method of making sulphuric acid is to treat sulphur trioxide with water (cf. 190). The so-called " Kng- lish," or common, process consists in oxidizing sulphur dioxide in the presence of water. The oxidizing agent used is nitric acid. The sulphur dioxide is produced from sulphur, iron pyrites (FeS 2 ), or galena (PbS); its oxidation is carried out in large boxes lined with lead and called " the leaden cham- bers." Currents of air, of steam, and, occasionally, of nitric acid enter the leaden chambers along with the sulphur dioxide, and sulphuric acid is the result. The simplest equation is, S0 2 + H 2 O + H 3 S0 4 . MANUFACTURE OF SULPHURIC ACID. 175 Explanation. The nitric acid introduced ii>to the leaden chambers is reduced to nitric oxide, NO; a small amount of sulphur dioxide is thus oxidized directly by nitric acid. But the greater portion of the oxygen used comes from the air ; for FIG. 41. the nitric oxide takes up the oxygen of the air, forming nitro- gen dioxide (c/. 170), and then gives up the oxygen to the sulphur dioxide. All of these facts are represented by the following equations : (1) (2) (3) (4) Repetition of (2). Theoretically, a very small amount of nitric acid ought to be able to oxidize an indefinitely large amount of sulphur dioxide, but in practice some of the nitrogen oxides are lost ; hence nitric acid must be added from 176 SULPHUR AND ITS COMPOUNDS. time to time to the mixture in the leaden chambers. The greater portion of the nitrogen oxides is prevented from escaping by being made to pass through towers of dilute sulphuric acid, which absorbs them. They are thus compelled to perform the work of oxidizing sul- phur dioxide over and over again. The equations given above are only incomplete representa- tions of the reactions taking place in the leaden chambers. It is probable that the nitric oxide reacts with the steam, oxygen and sulphur dioxide present in the leaden chambers, giving a substance called nitrosyl sulphuric acid. The equa- tion is, 2 SO 2 + 2 NO + 3 O -f IT./) = -2 XO. TISO 4 . The nitrosyl sulphuric acid is a solid substance. It is known technically as " chamber crystals." It is readily decomposed by the excess of steam, giving sulphuric acid and nitrogen trioxide (N" O 3 ), as is shown by the equation, 2 NO. HSO 4 + H 2 O = 2 II 2 SO 4 -f N,O 3 - Apparatus for demonstrating the nitrosyl sulphuric acid manufacture is shown in Fig. 41. 185. Purification of Sulphuric Acid. The sul- phuric acid obtained in the leaden chambers contains about 40^o of water; it is therefore concentrated by evaporation. The evaporation is carried out in leaden pans until the acid is concentrated enough to attack the lead. When this is the case, further evapora- tion is carried out in cast-iron pans until an acid con- taining about 13J& of water is obtained. This is REDUCTION OF SULPHURIC ACID. 177 the "crude" sulphuric acid of commerce. It is very impure. If the acid is to be concentrated still further, the evaporation must be carried out in vessels of glass, porcelain, or platinum. The pure acid is made by distilling the crude product in stills of platinum lined with gold. There is thus obtained an oil boiling at 338 C. and containing only 1.5% water. Its specific gravity is 1.S54 at C. The anhydrous acid (approximately 100% sulphuric acid) is made only in very small amounts. 1 86. Properties. Sulphuric acid is a thick, oily, colorless liquid. When it is diluted with water, much heat is evolved, so much, indeed, that the water some- times boils. To avoid spattering of the hot liquid we pour the concentrated acid in a small stream into the water, not the water into the acid. Sulphuric acid forms several hydrates with water, ^the two most important being those represented by the formulas, H 2 SO 4 . H 2 O and H 2 SO 4 . 2 H 3 O. Because of the tendency of sulphuric acid to take up water, it is used as a drying agent. The dehydrating power of sulphuric acid also accounts for the fact that organic matter, e. ? , SO 2 + H 2 O), 184 SULPHUR AND ITS COMPOUNDS. By this reaction a sulphite may readily be distinguished from a sulphate. Sulphites, like sulphurous acid, oxidize readily in the air. 194. Thiosulphates. Thiosulphuric acid, H 2 S 2 O 3 , is sulphuric acid with one-fourth of its oxygen replaced by sulphur (" thion " = sulphur) ; its salts are called thio sulphates. The most important thiosulphate is the sodium salt, Na 9 S 2 O 3 . This is made by boiling a solution of sodium sulphite with sulphur. Na 2 SO 3 + S = Na 2 S 2 O 3 . This reaction corresponds to the oxidation of sulphites to sulphates () impurities. The impurities are chiefly nitrogen find. carbon dioxide. The combustible gases are of two kinds : - (1) Those that burn without giving light, and (2) those that burn with luminous flames. The non-illuminating gases are hydrogen, methane, and carbon monoxide ; they make up about ninety per cent by volume of the gas. The illuminating gases are the so-called " heavy hydrocarbons," ethane, propane, 3 H 8 , butane, C 4 H 10 , etc., and, usually, small amounts of ethylene and acetylene. Two general processes are employed in making illumi- nating gas : (1) The distillation of soft coal. (2) The " water-gas " process. 223. Illuminating Gas by Distillation of Coal. - The old process of making gas is carried out as shown in Fig. 49. Soft coal in the retorts C (there are usually several retorts, one over the other) is heated by the fire A to the temperature ILL UMINA TING GAS B Y DIti TILL A TION OF COAL. 209 of decomposition. The volatile products pass off through the pipe T into the " hydraulic main " B. The hydraulic main contains water, which condenses much tar, etc., from the gas. From B the gas passes through the " condensers " D, which stand over water ; here the gas is cooled, and more of the tarry products are condensed. FIG. 49. From the condensers the gas passes through the coke towers, or " scrubbers " O ; into these water is sprayed, and the illuminating gas is thereby freed from soluble gases, such as ammonia and hydrogen sulphide. From the " scrubbers " the gas passes into the " purifiers " M. The purifiers are large boxes containing trays of slaked lime ; this substance removes carbon dioxide and traces of hydrogen sulphide. The gas then enters the gas holder G ; thence it is distributed to the community through the service pipe S'. Many other valuable products besides illuminating 210 CARBON AND ITS COMPOUNDS. gas are obtained by the distillation of coal. The am- moniacal liquors of the condensers arid coke towers are the sources of ammonium compounds ( region of unburned gases. These parts are shown in vertical section in Fig. 51. -B X is the region of unburned gases. B is the luminous zone. It contains solid par- ticles of carbon in a state of combustion. B' is the ruddy tip of the luminous zone. A is the outer mantle of the flame. Being non- luminous it is obscured by the light of B, except at the bottom, where it forms a blue, cup-shaped region. In addition to the parts just named, an impor- tant region is believed to exist about the region X. This zone is designated C in the idealized section of a candle flame (Fig. 52). Being non-luminous, the re- gion C is obscured by the light of B. -A FIG. 51. 216 FLAMES. Whether a flame is to be luminous or non-luminous depends upon the condition of affairs in the region O. What takes place in a candle flame is probably as follows : The vaporized paraffin (wax) of the region X (Fig. 52) burns in part in the zone (7, producing enough heat to decompose some of the paraf- fin vapor into hydrogen, certain hydrocar- bons (especially acetylene), and solid carbon. These substances burn further in the region B, the carbon burning, as usual, with a bright -X glow, and thus causing the luminosity of this -A re gi n ' I 1 ' 1 A the gases and the carbon escap- ing unburned through B are more or less completely burned. FIG. 52. 230. Non-Luminous Flames. The de- composition of the combustible material of region X (Fig. 52) into acetylene, carbon, etc., in the region C requires a definite degree of temperature ; hence, if the temperature of is sufficiently lowered, the flame becomes non-luminous. This is exactly what takes place in practice ; for if large quantities of a cold diluting gas, e. g., air, carbon dioxide, or nitrogen, are introduced into X, thus cooling C, the luminosity of the flame is destroyed. If, however, the diluting gases are first heated, the non-luminous flame becomes lumi- nous. SIMPLE AND COMPLEX FLAMES. 217 In the non-luminous flame the region C may readily be distinguished by its lightrblue color. The explanation of non-luminosity in flames, just' given, contains the theory of the Bunsen burner (Fig. 53). When the holes- at the base of the burner are open, the gas which rushes past the holes draws in currents of air, and the flame is non- luminous; when the holes are closed the flame is luminous. Carbon dioxide, nitro- gen, etc., give the same result as air. Air I Air The zones of combustion in the Bunsen flame are as follows (Fig. 54):- FIG. 53. .XT is the region of unburned gas, as in the candle flame. G is the light-blue, inner cone surrounding X. B is the non-luminous, dark cone. In the candle flame this is luminous. A. A is the purple, outer mantle. In the candle this is the faintly luminous liuJlo surrounding the flame. 231. Simple and Complex Flames. The C simplest flames are those having only one cone of combustion. Illustrations are the flames of .A. carbon monoxide (y experiment. If AVC call the amount of heat necessary to raise the temperature of one gram of water from C. to 1 C. a ymm-centi- t/rtide (g. c.) heat unit, we may write the equation for the union of hydrogen and oxygen as follows : H 2 -f O = IT,() -f 08,400 g. c. heat units. This equation shows not only that 2 grams of hydro- gen and 16 grams of oxygen unite to form 18 grams of water, but also that by their union 68,400 heat units are liberated. Similarly, for the union of 12 grams of carbon and 32 grams of oxygen we may write, C + O 2 = flQ 3 -f- 97,000 g. c. heat units. For the union of 1 gram of hydrogen and 35.5 grams of chlorine the equation is, H -j- Cl = HC1 + 22,000 g. c. heat units. The difference between the energy (calculated as heat) possessed by the elements which united to form a com- pound and that possessed by the compound itself is called the heat of formation of the compound. The heat of decomposition of a compound is numeri- cally equal to the heat of formation ; that is to say, the quantity of heat necessary to separate a compound into its elements is just as great as that evolved when the compound was formed from its elements. HEAT OF FORMATION EVOLVED IN STAGES. 221 235. Positive and Negative Heat of Formation. - The heat of formation of water, of carbon dioxide, and of hydrochloric acid is positive (-f-) heat being evolved when these compounds are formed; many cases, how- ever, exist in which heat is not evolved, but absorbed in the formation of a compound from its elements. In such cases the heat of formation is negative ( ). An illustration is the case of carbon disulphide, a substance which is produced (c/. 183) by passing sulphur vapor over hot charcoal. Carbon burns in sulphur with absorption of heat. The quantity of heat rendered potential by the union of 12 grains of carbon and 64 grams of sulphur is shown by the equation, C -f 2 S = CS 2 19,600 g. c. heat units. Similarly, hydrogen and-iodine unite with absorption of heat. H -f I = HI 6,100 g. c. heat units. A compound with a negative heat of formation is in a state of tension, or of unstable equilibrium ; for it is possessed of more energy than its constituent elements. When such a compound is decomposed, energy is evolved. 236. Heat of Formation Evolved in Stages. Just as it makes no difference in the total quantity of heat given out whether a given mass of a combustible burns slowly or rapidly (cf. 28), so the total amount of heat evolved (or absorbed) in the formation of a compound is the same whether the compound is formed in one or in several stages. 222 HEAT OF FORMATION AND DECOMPOSITION. Thus, the heat of formation of calcium carbonate, CaCO 3 , is equal to the sum of the heats of formation of calcium oxide, CaO, and of carbon dioxide, CO 2 , plua the heat evolved when cal- cium oxide and carbon dioxide unite to form calcium carbon- ate, CHAPTER XVI. MOLECULES AND ATOMS. 237. Law of Multiple Proportions. We have already learned (). All of the cases named illustrate the Law of Multi- ple Proportions, which may be stated in its simplest form as follows : If two elements form several compounds with each other, the different masses of one element which combine ivith a fixed mitxs of the other /><'<() CO 2 , tells us that one vol- ume of oxygen disappears in the formation of one volume of carbon dioxide ; but it does not tell us the volume of the car- bon that unites with one volume of oxygen, since we cannot experiment with gaseous carbon. 259. Valence. The student must have noticed from the formulas previously studied that atoms differ greatly in their power of combining with other atoms. Thus, the formulas of the compounds of hydrogen show interesting differences ; for, while in the case of hydro- chloric acid one atom of chlorine unites with one of hydrogen, in the case of water (H 2 O) one atom of oxygen holds two of hydrogen. The combining powers of the nitrogen and the carbon atom are still greater ; for in the molecule of ammonia (NH 8 ) one nitrogen atom holds three hydrogen atoms ; and in the molecule of marsh gas (CH 4 ) one carbon atom holds four atoms of hydrogen. This power of the atoms to unite with different num- bers of other atoms is called valence. An element like chlorine, whose atom can hold only one atom of hydro- gen, is said to have a valence of one, or to be univa- 248 MOLECULES A.VD ATOMS. lent. The hydrogen atom is always considered univa- lent. An element like oxygen, whose atom can hold two hydrogen atoms, is said to have a valence of two, or to be bivalent. One like nitrogen is, therefore, trivalent J and one like carbon, quadrivalent. Valence is not only the power of combining with, but also of replacing, different numbers of atoms. Thus, i the formula of potassium sulphate, ^ SO 4 , is derived TT from that of sulphuric acid, , j SO 4 , by replacing two hydrogen atoms by two of potassium. Potassium has, therefore, a valence of one. In the case of zinc nitrate, Zn XT r 3 ' one atom of zinc has re- JMJg, placed two hydrogen atoms (in 2 HNO 8 ) ; hence zinc is bivalent. Similarly, iron is trivalent in ferric phosphate, FePO 4 ; for the formula of phosphoric acid is H 8 PO 4 . Aluminum, also, is trivalent in aluminum sulphate, A1 2 (SO 4 ) 3 ; for two aluminum atoms, each with a valence of three, replace six hydrogen atoms (in 3 H 2 SO 4 ). The valence of an element is often represented by small Roman figures placed a little above and to the right of the symbol of the element. Thus, A I 111 means trivalent aluminum ; Hg 1 , univalent mercury ; and Pt IV , quadrivalent platinum. 260. Different Formula Types Based on Valence. A bivalent element, like oxygen, unites with two GRAPHIC FORMULAS. 249 atoms of a univalent element, but with only one of an- other bivalent element. Thus, calcium chloride has the formula CaCl 2 ; but calcium oxide has the formula CaO. When a bivalent element unites with a trivalent ele- ment, two atoms of the trivalent element generally require three of the bivalent one. This is shown in the formulas A1 2 O 3 for aluminum oxide, As 2 S 3 for arsenious sulphide, and Mg 8 N 2 for magnesium nitride. When quadrivalent atoms, like those of carbon and silicon, unite with univalent atoms, four of the univa- lent atoms are required, as in CH 4 and SiCl 4 ; when they unite with bivalent atoms, two of the latter are usually required, as in CS 2 and SiO 9 . The valence of an element is not, however, fixed ; for car- bon forms the compound CO, in which its valence is undoubt- edly two. So, also, nitrogen is trivalent in the compounds NTI 3 and N 2 O 3 , but quinquivalent in NH 4 C1 and N 2 O 5 . 261. Graphic Formulas. By means of the idea of valence, we may represent the relation of atoms to one another in molecules. When the molecule consists of two atoms, only one arrangement is possible, viz., the atoms are joined directly. Thus in hydrochloric acid and in molecular hydrogen we have simple union. If we repre- sent the combining power of the elements by lines (called bonds), we may write the formula for hydro- chloric acid graphically H Cl and that of hydrogen H H. The single line shows that the valence of each atom in the molecule is one, 250 MOLECULES AND ATOMS. Everything goes to show that the two hydrogen atoms in the water molecule are not united to each other, but to oxygen. Similarly it is believed that the hydrogen atoms of the am- monia molecule are all united to nitrogen, and those of the marsh-gas molecule to carbon. We may represent these facts in the formulas, H ^11 | II O II ; ; and II C II. Formulas like those just given are called graphic 01 structural formulas. We may represent an element in its nascent state by the symbol with free valence bonds. Thus, O represents nascent or atomic oxygen ; H , nascent hydrogen. 262. Isomerism. Graphic formulas enable us to represent differences between compounds which cannot be distinguished by the ordinary formulas. Methyl ether and ethyl alcohol, for example, are both repre- sented by the formula C 9 H 6 O ; these substances are so different, however, that no one would mistake one for the other. Thus, methyl ether boils at 24 C., at or- dinary pressure, while ethyl alcohol boils at -)-78 C. Such compounds are said to be isomeric with each other. By the use of graphic formulas, the difference in character between these isomeric substances is readily understood. Thus, the graphic formula for methyl ether is, ALLOTROPISM. 251 H II I I H C O C H; I I II H while that of alcohol is, II II I I H C C O H. I I II H According to these formulas, all the hydrogen atoms of methyl ether have the same relation to the remainder of the molecule and should behave in the same way with reagents; while in the case of ethyl alcohol one hydrogen atom the one bound to oxygen should be different from the other five. This is actually the case; for the atom of hydrogen bound to oxygen is the only one of the six that can be re- placed by sodium and other metals. 263. Allotropism. Just as there are compounds having the same chemical composition which are yet very unlike in their properties, so there are elements existing in forms so different that they might easily be sup- posed to be entirely different substances. In the pre- ceding section we have called the compounds isomers ; the different forms of the same element are called allo- tropic forms of the element. The existence of an element in different forms is called allotropism. Car- bon and sulphur, as we have already learned, exist in several allotropic forms ; the same is true of oxygen, phosphorus, silicon, boron, etc. 252 MOLECULES AND ATOMS. Allotropism is probably due to different causes, such as different arrangements of the atoms in the molecule, or differ- ent numbers of atoms in the molecule. In many cases an allotropic form is only temporary. This is true of the plastic, or amorphous, modification of sulphur (cf. 175), which changes into the ordinary form with evolution of heat. Plastic sulphur thus represents a condition of unstable equilibrium, like a compound which has a negative heat of for- mation (cf. 235). 264. Exercises. 1. (a) A liter of a certain gaseous substance weighs approxi- mately 1.966 grams at standard conditions ; what is its molecu- lar mass ? (See 242.) (6) If -ft- of this substance is carbon and -ft oxygen, what is its formula ? (See 252.) 2. (a) Two hundred c.c. of a gas weigh 0.3932 grams at standard conditions ; what is the molecular mass ? (6) Analysis shows that the gas is composed of nitrogen, 63.64%, and oxygen, 36.36% ; what is the formula? 3. The oxide of magnesium is composed of magnesium, 60%, oxygen, 40%; what is its simplest formula? 4. A chloride of phosphorus has the composition, phos- phorus 22.55%, chlorine 77.45% ; find its simplest formula. 5. What is the approximate atomic mass of platinum if its specific heat is about 0.033 ? 6. What volume of oxygen is used v up when 20 c.c. of acetylene burn in air ? What is the volume of carbon dioxide formed ? (Cf. 221 and 258.) 7. Write the molecular equation for the combustion of pen- tane, C 5 H 10 , in oxygen, if the products are carbon dioxide and water. What volume of oxygen is used up when 50 c.c. of pentane burn? EXERCISES. 253 What volume of carbon dioxide is produced ? 8. Knowing that the valence of an element x is 1, write the simplest formula for its sulphate, its carbonate, and its nitrate. Q. Write the simplest formulas of the chloride and sulphite of an aliment whose valence symbol is Si ri . CHAPTER XVII. FLUORINE, BROMINE, IODINE, AND THEIR COMPOUNDS. 265. Halogens. The elements fluorine, chlorine, bromine, and iodine are called " the halogens," from hah, Greek for "salt," and the suffix gen, meaning " a constituent of," as in " hydrogen," etc. Fluorine and chlorine are gaseous at the ordinary tempera- ture ; bromine is a liquid boiling at about 50 C. ; while iodine is a black solid which gives off, even at ordinary temperatures, a beautiful, violet vapor. 266. Fluorine. The element fluorine was known in its compounds long before it was obtained in the free condition. The most common of its compounds is cal- cium fluoride, or fluorspar, CaFl 2 . Fluorspar derives its name frqmfliw, Latin for "to flow," and spar, meaning " a rock." The name is applied to this substance owing to the use of fluorspar as 'A flux in metallurgy. A flux is an easily fusible substance added to the mixture of an ore and a reducing agent to promote fusion of the mix- ture. The substance resulting from the union of the flux with the impurities present is usually called " slag." Another important natural fluorine compound is cryo- lite. This is a double fluoride of aluminum and sodium; its formula is A1FL. 3 NaFl or Na Q AlFL. o o o 254 HYDROFLVOHIC ACID. 255 Fluorine cannot be prepared from either of these compounds directly, but has been made (1886) by the electrolysis of anhydrous hydrofluoric acid, HFL 2 HF1 = H 2 + F1 2 . The operation may be carried out in copper or plati- num apparatus, but not in glass, since hydrofluoric acid attacks glass energetically. 267. Properties of Fluorine. Fluorine is a yellow gas, about one and two-fifths times as heavy as air. It acts upon water with violence, according to the equation, 2 H 2 O + 2 Fig 4 HF1 + O 2 . The oxygen formed always contains some ozone (c/. 287). Fluorine unites with hydrogen explosively, even in the dark (cf. Chlorine, 84), to form hydrofluoric acid. It forms no compounds with oxygen, so far as known. As commonly prepared, fluorine acts upon glass, but this is due to the fact that a small amount of hydro- fluoric acid is present in the fluorine. Fluorine acts readily upon silicon and antimony, forming the corresponding fluorides, SiFl 4 and SbFl 3 . Gaseous fluorine has been condensed to a liquid boiling at 187 C. at ordinary pressure. 268. Hydrofluoric Acid. Hydrofluoric acid is com- monly prepared by heating calcium fluoride with con- centrated sulphuric acid. The equation is, - CaFl 2 + H 2 S0 4 = CaSO 4 + 2 HFL 256 FLUORINE, BROMINE, IODINE. Anhydrous hydrofluoric acid is a liquid boiling at about 19 C. Both the vapor of this liquid and its solution in water are very poisonous. The aqueous solution reacts with almost all the metals, forming fluorides and hydrogen ; and decomposes the oxides, forming fluorides and water. Silicon dioxide (quartz, sand, etc.) gives with hydro- fluoric acid silicon tetrafluoride (SiFl 4 ) and water, ac- cording to the equation, SiO 2 + 4 HF1 = SiFl 4 + 2 H 2 O. Silicon tetrafluoride is a gas. Glass a mixture of silicates, i.e., salts of silicic acid, H 2 SiO 3 is acted upon by hydrofluoric acid as silicon dioxide is. Thus, we may represent the action of calcium silicate, CaSiO 3 , with the acid by the equa- tion, CaSiO 3 + 6 HF1 = CaFl 2 + SiFl 4 + 3 H 2 O. Hence, when glass is treated with hydrofluoric acid, the sili- con present in the glass escapes as SiFl 4 , leaving a depression in the glass. This fact is made use of in the operation of etch- ing glass. The glass is first covered with a thin layer of paraf- fin, and a design is drawn in the paraftin by means of a sharp point. When the exposed glass is wet with the solution of the acid (a swab of cotton attached to a stick may be used to ap- ply the solution), or is left in the vapor of the acid, the design is etched into the glass. Hydrofluoric acid is commonly kept in bottles of paper, covered inside and out with a thick layer of par- PREPARATION OF BROMINE. 257 affin. Vessels of lead, platinum, or rubber may also be used. Bromine. 269. Preparation of Bromine. Bromine is found in nature in the combined form, chiefly as bromides. The most common bromides are those of sodium (NaBr), of potassium (KBr), and of magnesium (MgBr 2 ). Bromides occur in sea-water and in connection with salt deposits. Bromine is prepared hy heating a bromide with man- ganese dioxide and dilute sulphuric acid. The bromine vapor evolved is condensed in cold receivers. With sodium bromide the equation is, Mn0 2 -|- 2 NaBr -f 3 H 2 SO 4 = MnSO 4 + 2 NaHSO 4 -f Br 2 -f 2H 2 0. This reaction is like that used in making chlorine (cf. 81) from common salt, manganese dioxide, and sulphuric acid. The reaction takes place in at least two stages : (1) The sulphuric acid and sodium bromide give sodium hy- drogen sulphate and hydrobromic acid, according to the equa- tion, NaBr + H 2 S0 4 NaHSO 4 + HBr ; and (2) The hydrobromic acid and the manganese dioxide react to give manganous bromide (MnBr 2 ). MnO 2 + 4 HBr = MnBr 2 + Br 2 -f 2 H 2 O. The manganous bromide and sulphuric acid then give rise to manganous sulphate (see above) and more hydrobromic acid. MnBr 2 + H 2 SO 4 > MnSO 4 + 2 HBr. 258 FLUORIXE, BROMINE, IODINE. Bromine may also be prepared by conducting the proper amount of chlorine into the solution of a brom- ide. With magnesium bromide the equation is, MgBr, + C1 3 > Mg('l 2 + Br r 270. Properties of Bromine. Bromine is a brown liquid about 3.2 times as heavy as water. Its vapor has an odor much like that of chlorine, and affects the eyes. Bromine boils at about 59 C. The density of bromine vapor shows that the mole- cule is diatomic ; its formula is, therefore, Br 2 . At about 1000 C. the molecule begins to dissociate into molecules containing only one atom each (cf. 45). Bromine dissolves in water, carbon disulphide, and other solvents. The aqueous solution is called " brom- ine water." In the presence of some substance capable of taking up oxygen, bromine reacts with water energetically, according to the equation, H 2 + Br 2 > 2 HBr + ( O -). By O we mean nascent oxygen (cf. 261). Bromine water is thus a good oxidizing agent. The same action goes on more slowly when no oxidizable substance is present ; bromine water thus becomes converted into a dilute solution of hydrobromic acid, HBr. Bromine is less active than chlorine, but, like chlor- ine, it unites with hydrogen and with metals to form bromides. PROPERTIES Of HYDROBROMIC ACID. 259 271. Hydrobromic Acid. Hydrobromic acid cannot be made in a pure state by treating a bromide with con- centrated sulphuric acid, for the reason that some of the hydrobromic acid formed breaks up into hydrogen and bromine, and the nascent hydrogen reduces the sulphuric acid. The products of the action are thus bromine and sulphurous acid, as well as hydrobromic acid. These facts are represented in the equations, (1) NaBr 4- H 2 SO 4 > NaHSO 4 -f HBr ; (2) 2 HBr + H 2 S0 4 > H 2 O + Br 2 + H 2 SO 3 . The method commonly used to prepare hydrobromic acid is to treat red phosphorus with bromine in the presence of water. The phosphorus and bromine first unite to form phosphorus tribromide, PBr 3 ; but this substance is decomposed at once by the water to form phosphorous acid and hydrobromic acid. The equations are, (1) 2P+3Br 2 :=2PBr 3 ; (2) PBr 3 + 3 H 2 = H 3 PO 3 + 3 HBr. The phosphorous acid, being non-volatile, remains behind ; while the gaseous hydrobromic acid passes off. The hydrobromic acid may be freed from bromine vapor by passing it through a U-tube containing moist red phosphorus. 272. Properties of Hydrobromic Acid. Hydro- bromic acid gas is like hydrochloric acid gas. It fumes in the air and dissolves readily in water. Its concen- 260 FLUORINE, BROMINE, IODINE. trated aqueous solution has a specific gravity of almost 1.8, and contains 82J& by weight of the acid. Hydrobromic acid begins to dissociate into its elements (c/. 45) at about 800 C.; it is, therefore, much less stable than hydrochloric acid, which begins to dissociate at about 1500 C. Iodine. 273. Occurrence and Preparation of Iodine. The chief source of iodine until recently was the ashes of certain sea-plants which absorb iodine compounds from sea-water. At the present time the element is obtained largely from the Chile saltpeter deposits. In these deposits the iodine is found chiefly as sodium iodate, NaI0 3 . Iodine may be set free from an iodide in just the same way that chlorine and bromine are set free from chlorides and bromides respectively, viz., by heating it with a mixture of manganese dioxide and dilute sulphuric acid. A representative equation is, MnO 2 + 2 Nal + 3 H 2 SO 4 = MnSO 4 + 2 NaHSO 4 + 2H 2 + I 2 . The stages in which the reaction takes place are partly represented by the equations, (1) Kal + H 2 SO 4 = HI -f NaHSO 4 ; (2) Mn0 2 + 4 HI = MnI 2 + 2 H 2 O + 2 I; (3) MnI 2 + H 2 S0 4 = MnS0 4 + 2 HI. Iodine may also be set free from an iodide by means of chlorine or bromine (cf. 269). HYDRIODIC ACID. 261 2 Nal + Br 2 2 NaBr +21. 2 Nal + C1 2 " 2 NaCl + 2 I. 274. Properties of Iodine. Iodine is, ordinarily, an almost black solid, melting at 114 C. and boiling at about 184 C. Its vapor has a beautiful, violet color; it is about 8.7 times as heavy as air. Iodine is very soluble in carbon disulphide and in ether, but only slightly soluble in water. It is less active than chlorine or bromine. It stains the skin brown, and imparts an intensely blue color to starch paste. Iodine sublimes (ef. 149) when heated. Up to about 600 C. the molecule of iodine vapor consists of two atoms ; above this temperature, dissociation takes place. At about 1500 C. only monatomic iodine molecules exist. 275. Hydriodic Acid. Hydriodic acid is still more unstable than hydrobromic acid. It cannot be made in a pure condition by treating an iodide with concentrated . sulphuric acid for the reason that the hydriodic acid reduces the sulphuric acid. The reduction goes not only to sulphurous acid, as in the case of hydrobromic acid, but in part even to hydrogen sulphide. The equa- tions representing this are, (1) KI + H 2 SO 4 = KHSO 4 + HI ; (2) 2 HI + H 2 S0 4 = H 2 S0 3 + H O + 21; (2a) 8 HI + H 2 S0 4 = H 2 S + 4 H 2 O + 81. Hydriodic acid gas may be made by allowing red phosphorus and iodine to react in the presence of water. FLUORINE, BROMINE, IODINE. Phosphorus tri-iodide is first formed, but is decomposed at once according to the equation, PI S + 3 H 2 = H 8 P0 8 + 3 HI (cf. 271). An aqueous solution of hydriodic acid may best be prepared by making use of a property common to chlorine, bromine, and iodine, viz., the ability of each of these substances to decompose hydrogen sulphide. The sulphur formed, being insoluble, is precipitated ; hence the reaction goes on to completion (cf. 180). The equation in the case of iodine is, 2 I -f H 2 S > 2 HI + S. The hydrogen sulphide is conducted into a mixture of iodine and water until the iodine disappears. The sulphur is then filtered off, and the filtrate distilled. After the water has passed off, a heavy liquid is obtained, which boils at 126 0. This is about 57% hydriodic acid. 276. Properties of Hydriodic Acid. Hydriodic acid gas is about 4.4 times as heavy as air. Like hy- drochloric and hydrobromic acids, it is very soluble in water. One cubic centimeter of water at 10 C. and standard pressure dissolves about 450 c.c. of the gas. Hydrogen and iodine can be made to unite under appropriate conditions. When uniting they do not evolve heat, but absorb it. This accounts for the fact that hydriodic acid is so unstable (cf. 235). The dissociation of hydriodic acid is like that of steam (cf. 45). At any temperature above the point COMPOUNDS OF THE HALOGENS WITH OXYGEN. 263 at which dissociation begins, the decomposition of hydri- odic acid into hydrogen and iodine goes on side by side with recombination of hydrogen and iodine to form hy- driodic acid. The condition of equilibrium is reached when as many molecules of hydriodic acid are formed in a given time as are decomposed in the same time. We may represent this condition of equilibrium by the equation, 2 HI jn H 2 + I 2 . Such an equation is called a "balanced" or "equilib- rium " equation. The arrows indicate that the reac- tion goes in both directions. If either product of dissociation is removed from the " sphere of action," the dissociation goes on to comple- tion. Thus, if silver is placed in hydriodic acid solu- tion, it unites with the iodine as rapidly as iodine is formed. Hence hydrogen is set free. 2 Ag + 2 HI 2 Agl + H 2 . Because of its ready dissociation, hydriodic acid acts as a powerful reducing agent. Oxygen, or any oxidizing agent, gives with it iodine and water. 4 HI + 2 2 H 2 + 4 I. 277. Compounds of the Halogens with Oxygen. Only three halogen oxides have been actually made ; although more especially one oxide of bromine are suspected to be capable of existing. The three oxides definitely known are : 264 FLUORINE, BE MINE, IODINE. Chlorine monoxide, C1 2 O ; Chlorine dioxide, C1O 2 or C1 2 O 4 ; Iodine pentoxide, I 2 O & . Chlorine monoxide is an unstable liquid which boils at -)-5 C. It explodes with violence. When one volume of chlorine monoxide is carefully decomposed, it gives half a volume of oxygen and one volume of chlorine. Chlorine dioxid^ (C1O 2 ) and tetroxide (C1 2 O 4 ) corre- spond to the nitrogen oxides NO 2 and N 2 O 4 . It is formed with violent explosion when concentrated sul- phuric acid acts upon potassium chlorate. The prepara- tion of this substance should not be attempted without precise directions and extraordinary precautions. Chlorine dioxide is a reddish-brown liquid, boiling at about 10 C. Iodine pentoxide is the only oxide of iodine known at present. It is a white, stable powder; with water it gives iodic acid, I 2 6 + H 2 = 2 HI0 3 . Iodine pentoxide is thus the anhydride of iodic acid. 278. Compounds of the Halogens with Oxygen and Hydrogen. Compounds of the halogens with oxygen and hydrogen are called oxygen acids, or oxy-acids, of the halogens. They are more numerous than com- pounds with oxygen alone ; for they include at least three chlorine acids, three bromine acids, and two iodine HYPOCHLOROUS ACID. 265 acids. The formulas of these acids appear in the fol- lowing table : HC10. IIBrO. IiriO ( known only | 1U2 I in its salts J HC10 3 . HBr0 3 . HIO 3 . HC1O 4 . HBrO 4 . HIO 4 . 279. Hypochlorous Acid, H-0-C1. Hypochlorous acid is present in a solution of chlorine in water ( 2 HC1 + ( O ). In the presence of sunlight a concentrated solution of chlor- ine in water gives off oxygen (c/. 85). Salts of hypochlorous acid, i. e., hypochlorites, are formed by passing chlorine into dilute solutions of hydroxides of the metals. The amount of chlorine used must be less than is required to saturate the hy- droxide. With potassium hydroxide the equation is, 2 KOH + C1 2 KOC1 + KC1 + H 2 O. 266 FLUORINE, BROMINE, IODINE. When chlorine acts upon powdered slaked lime, Ca(OH) 2 , Ueaching powder (Ca^p, \ is produced. C1 /OC1 280. Chlorous Acid, H-0-C1=0. Chlorous acid does not exist free. Its potassium salt is produced, iu solution, when an aqueous solution of chlorine dioxide, C1O 2 , is treated with potassium hydroxide. 281. Chloric Acid, H-O-Cl^Q. Chloric acid is known only in its aqueous solution ; this may be con- centrated until it contains 40^ of chloric acid. Chloric acid is a powerful oxidizing agent. Chlorates are formed by conducting chlorine into hot, concentrated solutions of alkalies to complete saturation. With potassium hydroxide the equation is, 6 KOH + 3 C1 2 - KC1O 3 + 5 KC1 + 3 H 2 O. The chlorate is separated from the chloride by recrystalliza- tion from hot water. The potassium chlorate, being much less soluble than the other, separates out first. Potassium chlorate has already been used to produce oxygen ( 19). The equations representing the action of hydrochloric acid upon potassium chlorate (82) are, in part, (1) KC1O 3 + HC1 - KC1 + HC1O 3 ; (2) HC10 8 -j- 5 HC1 - > 3 H 2 + 3 Cl a . COMPOUNDS OF BROMINE WITH OXYGEN, ETC. 267 ^ 282. Perchloric Acid, H-0-C1=0. Perchloric acid ^0 is a colorless, explosive liquid about 1.8 times as heavy as water. Its salts, the perchlorates, are produced when the chlorates are partly decomposed by heat. In the decomposition of potassium chlorate (c/. 19), a point is soon reached at which a considerable increase of temperature is needed to continue the evolution of oxygen. The amount of oxygen evolved up to this point is only one-third of the quantity present in potassium chlorate. If we stop at this stage, we obtain a mixture of potassium perchlorate and potassium chlor- ide, as shown in the equation, 2 KC1O 3 = KC1O 4 + KC1 + O 2 . The potassium chloride is much more soluble than the per- chlorate; hence these substances may be separated by recrystal- lization from water. 283. Compounds of Bromine with Oxygen and Hy- drogen. No compounds of bromine and oxygen have been prepared in a pure condition. With oxygen and hydrogen bromine forms kypobromous acid, bromic acia, and perbromic acid. The graphic formulas of these acids are like those of the corresponding chlorine compounds. Hypobromites are formed when cold, dilute alkalies are treated with bromine, but with less than is required for saturation. 2 KOH + 2 Br KOBr + H 2 O + KB:-. 268 FLUORINE, BROMINE, IODINE. The hypobromites, like the hypochlorites, are oxidizing agents. Bromates are formed when hot alkalies are saturated with bromine (ef. chlorates, 281). Potassium bromate is a white, crystalline solid like potassium chlorate. Heat decomposes it into potassium bromide and oxygen. 2KBrO 3 =2KBr+3 O 2 . 284. Compounds of Iodine with Oxygen and Hydro- gen. Two oxy-acids of iodine are known ; they are iodic and periodic acids. lodic acid (HIO 3 or, graphically, H-O-I^^ J is formed by oxidizing iodine with concentrated nitric acid. It is a crystalline solid. At 170 C. it breaks up into iodine pentoxide (I 2 O 5 ) and water. lodates are formed by adding iodine to hot, concen- trated solutions of alkalies. With potassium hydroxide the equation is, - 6 KOH +61 - KIO 3 + 5 KT + 3 Periodic acid is HIO,, or H-O-I=O. Its salts are \> periodates, e. g., sodium periodate, NaIO 4 . 285. The Halogen Family. From the preceding pages it is evident that there is a great similarity in the properties of the elements fluorine, chlorine, bromine, and iodine. The close relation of these elements to one THE HALOGEN' FAMILY. 269 another is yet more marked if we consider that there is a gradation in the properties of these elements in the order of the atomic masses. Thus, the melting tempera- tures, the boiling temperatures, and the specific gravities of these elements rise from fluorine (atomic mass 19) to iodine (atomic mass 127). The intensity of the color of these elements also increases with the atomic mass : fluorine is greenish yellow ; chlorine, green ; bromine, brown; and iodine, black. The gradation observed in their physical properties is true also of their chemical properties. Thus, the elements of lower atomic mass can expel those of higher atomic mass from their soluble metal salts. The same gradation of properties is noticed in the compounds of the halogens. Thus, the specific gravity of the hydrogen compounds increases from hydrofluoric acid to hydriodic acid ; while the stability of these com- pounds decreases in the same order. In the case of the compounds of the halogens with oxygen, and with oxygen and hydrogen, the order of stability is re- versed, iodine forming the most stable ones, chlorine very un- stable ones, and fluorine none at all. The same gradation of color see^i in the elements themselves appears in many of their compounds. Thus, silver chloride is white; silver bromide, light yellow; silver iodide, bright yellow. Some of the above (and other) facts appear in the following table : 270 FLUORINE, BROMINE, IODINE. PROPERTIES. FLUORINE. CHLORINE. BROMINE. IODINE. Atomic Mass .... 19 35.5 80 127 Boiling Temperature 187 C. 33 +59' 4-184" Specific Gravity . . 1.15 (liquid) 1.5 (liquid) 3.2 (liquid) 5 (solid) Union with Hydro- gen takes place . In the dark at ordinary tempera- tures. In sunlight. At red heat. A trod heat, but incom- pletely. Pleat of formation of Hydrogen Com- 37.6 heat units. 22 8 6.1 Stability of Hydro- gen Compound . . Most stable. Decomposed at 1500* C. Decomposed at 800 C. Decomposed at 180 C. Stability of Oxygen Compound .... Forms none. Unstable. Forms none. Most stable. A group of elements related to one another like the halogens is called a " Natural Family of Elements/' Several other natural families exist, and will be referred to later. 286. Exercises. 1. How many grams of bromine can be made from 150 grams of potassium bromide by using manganese dioxide and dilute sulphuric acid ? 2. Calculate the per cent of hydrogen in hydriodic acid. In hydrobromic acid. 3. How would you separate a mixture of iodine and sand ? EXERCISES. 271 4. How could you distinguish, by chemical means, between a chloride, a bromide, and an iodide ? 5. How could you identify a fluoride, e. g. calcium fluoride ? 6. About how much would a liter of air weigh at -f-273 c. and 760 mm. pressure ? A liter of iodine vapor ? CHAPTER XVIII. OZONE AND HYDROGEN PEROXIDE. 287. Ozone. Oxygen which -has been exposed to the silent electric discharge possesses new properties. It has a peculiar odor and oxidizing powers not pos- sessed by ordinary oxygen. Thus, it oxidizes silver and mercury at once, whereas these metals are not acted upon by oxygen at ordinary temperatures. These new properties are due to the presence in the oxygen of another substance, called ozone. The name ozone is from the Greek ozein, to smell. The same substance is produced in almost every case of oxidation. An illustration of this is the slow oxidation of phosphorus. If moist phosphorus is placed in a covered vessel, the peculiar odor of ozone soon appears. Ozone is formed, also, in the electrolysis of water, and appears at the -f- elec- trode along with oxygen. When ozone is examined, it is found to contain nothing but oxygen ; it is, in fact, an allotropic form of oxygen (cf. 263). When oxygen changes into ozone there is a contraction of volume amounting to one- third of the volume of oxygen taken. Ozone is, there- fore, one and one-half times as heavy as oxygen. Its molecular mass is 48 instead of 32 ; hence the molecule of ozone contains three oxygen atoms, and is written O 3 . 272 PROPERTIES OF OZONE. 273 The peculiar instability of ozone is due to the fact that the change from oxygen to ozone is accompanied by absorption of heat. In the presence of a substance capable of taking up oxygen, the ozone molecule readily gives up an atom of oxy- gen, and thus reverts to molecular oxygen, O = O. 288. Properties of Ozone. As might be expected, the oxidizing power of ozone is very great. Moist phosphorus and sulphur are converted by it into phos- phoric and sulphuric acids, respectively ; and ammonia is at once oxidized to nitric acid. Organic coloring substances, e. g. indigo and litmus, are at once decolor- ized by ozone. The bleaching of fabrics on exposure to the air is probably due to the action of ozone pres- ent in the air. When ozone is heated, its molecule is decomposed, and ordinary oxygen results. The reversion of ozone to oxygen is accompanied by an expansion of volume just equal to the con- traction that takes place when oxygen changes into ozone. Ozone is readily absorbed by oil of turpentine ; hence the amount of ozone formed in a given volume of oxygen may be determined by exposing the ozonized oxygen to this substance. Only about six per cent of a given amount of oxygen can be converted into ozone, because the reverse change of ozone into oxygen soon produces a condition of equilibrium. The presence of ozone in ozonized air is readily de- tected by means of a mixture of potassium iodide and starch paste best upon a piece of filter paper. The ozone sets iodine free, probably according to the equa- tion, 2 KI + H 2 + 3 2 KOH + 2 1+ O 2 . 274 OZONE AND HYDROGEN PEROXIDE. 289. Hydrogen Peroxide. Closely related to ozone, and long confused with it, is hydrogen peroxide, H 2 O 2 . Hydrogen peroxide is a colorless liquid about one and one-half times as heavy as water ; it possesses remark- able oxidizing and reducing powers. A dilute solution of hydrogen peroxide may be made by adding barium peroxide, BaO 2 , to dilute hydrochloric acid. The equation is, BaO 2 + 2 HC1 BaCl 2 + H 2 O 2 . A somewhat better way is to treat a dilute solution of tar- taric acid with sodium peroxide, Na 2 O 2 . NajOj + H 2 C 4 H 4 6 = Na 2 C 4 H 4 6 +H 2 O 2 . Sodium peroxide is made (along with sodium monoxide, Na 2 O) by burning sodium in air or oxygen. When phosphorus is partly immersed in water, it acts upon the moist air to form both hydrogen peroxide and ozone. Both of these substances are formed, also, by holding a hydrogen flame against a piece of ice. Hydrogen peroxide is formed in the electrolysis of water, if oxygen is passed into the water at the negative ( ) electrode. The oxygen is reduced by the nascent hydrogen evolved at the electrode. Hydrogen peroxide is found in the air, and in all rain water and snow. 290. Properties of Hydrogen Peroxide. Hydrogen peroxide may be obtained almost pure by distilling a dilute aqueous solution of it at low pressure. The ap- PROPERTIES OF HYDROGEN PEROXIDE. 275 paratus for distilling at reduced pressure is essentially as shown in Fig. 57. FIG. 57. A distilling flask (A) is provided with a thermometer (B) and a capillary tube (C). The capillary tube allows a very small stream of air to be drawn through the apparatus. The distilling flask is connected air-tight with the condenser (D) and the receiver (E). The pressure, in millimeters of mercury, is indicated by the manometer (F). The air is exhausted at S by a water or mercury suction-pump. At 26 mm. pressure hydrogen peroxide boils at 6 9 C.; under the same pressure water boils at 27 C. ; hence the two substances can be separated readily. The aqueous solution of hydrogen peroxide has a bitter taste and produces white spots upon the skin. It is a powerful antiseptic. Hydrogen peroxide decomposes readily, especially in the presence of basic substances. The products are water, and oxygen in the nascent condition ; hence hydrogen peroxide is a powerful oxidizing agent. It 276 OZONE AND HYDROGEN PEROXIDE. decolorizes indigo, litmus, etc., as ozone does. It at once oxidizes hydrochloric acid to water and chlorine. Hydrogen peroxide acts also as a reducing agent with evolution of oxygen. Thus, it reduces mercuric oxide to mercury and sets oxygen free. HgO + H 2 2 Hg + 2 + H 2 0. One atom of each oxygen molecule (O 2 ) comes from the hydrogen peroxide, and the other from the oxide reduced. Potassium permanganate solution is at once decolorized by hydrogen peroxide, and potassium chromate and bichromate solutions are changed to a green color. All of these are reductions. Ozone and hydrogen peroxide reduce each other. 3 + H 2 2 > 2 + H 2 + 2 . Hydrogen peroxide, like ozone, decomposes with evolution of heat ; this fact accounts for its instability. The common test for the presence of hydrogen peroxide in a solution is to add to the solution in a test tuhe about two or three cubic centimeters of ether, and then one drop of potassium bichromate solution. When the test tube is shaken, the layer of ether is colored a beau- tiful blue, if hydrogen peroxide is present. 291. Composition of Hydrogen Peroxide. Hydro- gen peroxide is composed of 1.01 parts, by weight, of hydrogen to every 16 parts of oxygen. Its molecular mass is 34 ; hence its formula is H 2 O 2 . The graphic formula of hydrogen peroxide is H-O-Q-H. COMPOSITION OF HYDROGEN PEROXIDE. 277 The effect of the structure of the molecule, i. e., the way in which the atoms are united, upon the properties of substances, is admirably illustrated by the differences in the behavior of the two classes of dioxides. Thus, while calcium peroxide (CaO 2 ), sodium peroxide (Na 2 O 2 ), and barium peroxide (BaO 2 ) give hydrogen peroxide when treated with dilute acids, lead dioxide (PbO 2 ) and manganese dioxide (MnO 2 ) do not. The structure of all true peroxides is like that of hydrogen per- oxide. The graphic formula of sodium peroxide is, therefore, Ka O O N"a, and that of barium peroxide, O O. The \/ Ba graphic formula of the dioxides is different, that of manganese dioxide being, probably, Mn^, and that of lead dioxide CHAPTER XIX. THE NITROGEN FAMILY. PHOSPHORUS, ARSENIC, ANTIMONY, BISMUTH. A. Phosphorus. 292. Occurrence and Preparation of Phosphorus. Phosphorus is found in nature only in the combined form, chiefly in phosphates. The most abundant phos- phate is calcium phosphate, Ca 3 (PO 4 ) 2 . Calcium phosphate exists in the soil, and is taken up from it by plants. Animals consume phosphates in their food. The immediate source of most phosphorus is bone-ash, which contains about 60^ to 70^/ of its weight of calcium phosphate. The present process of making phosphorus is to heat calcium phosphate with charcoal and sand in the electric furnace. We can understand the chemical reactions involved in making phosphorus by considering them separately. (1) The calcium phosphate probably breaks up in the presence of the silica (sand) into quiclelinie and phosphoi-us pentoxide. Ca 3 (P0 4 ) 2 > 3 CaO + P 2 O 5 . (2) The silica and quicklime unite to -give calcium silicate. 3 CaO + 3 SiO 2 = 3 CaSiO 8 . (3) The charcoal reduces the phosphorus pentoxide to phosphorus. 278 PROPERTIES. 279 P 2 5 + 5 C 2 P+ 5 CO. Hence the complete equation is, Ca 3 (P0 4 ) 2 + 3 8i0 2 + 5 C = 3 CaSiO 3 + 5 CO + 2 P. The phosphorus escapes from the furnace as a vapor, and is collected under water. To purify it, it is redis- tilled and pressed in the liquid state (under water) through a bone-ash filter. The phosphorus thus ob- tained is a white and transparent solid. About three thousand tons of phosphorus are made every year. 293. Properties. Phosphorus, like sulphur, exists in several allotropic forms with widely differing proper- ties. Ordinary or yellow phosphorus has a specific gravity of about 1.8, melts at about 45 C., and boils at 287 C. It is insoluble in water, but dissolves readily in carbon disulphide, CS 2 . Phosphorus derives its name, which means " bearer of light " (cf. Latin, lucifer), from its property of phosphorescing, i. e., glowing, when exposed in the dark to moist air or other gases containing oxygen. This phenomenon is caused by slow combustion on the surface of the phosphorus. Ordinary phosphorus ignites in air at 40 C., and burns with a hot flame to phosphorus trioxide and pent oxide. 30 3 =2P 2 8 . 4 P + 5 2 = 2 P 2 6 . The spontaneous ignition of finely divided phosphorus has already been described (cf. 29). 280 THE NITROGEN FAMILY. Phosphorus unites readily with chlorine, bromine, and iodine even at the ordinary temperature. Two compounds of phosphorus and chlorine are possible, viz., the trichloride, PC1 3 , and the pentachloride, PC1 5 . Phosphorus trichloride is a liquid ; phosphorus penta- chloride, a crystalline solid. 2 P -f 3 C1 2 = 2 PC1 8 . 2 P -f 5 C1 2 = 2 PC1 6 . Yellow pliosphorus is very poisonous. 294. Red Phosphorus. A great difference exists between yellow phosphorus and the red modification. Red phosphorus is a reddish powder 2.2 times as heavy as water, infusible at red heat, unable to phosphoresce, insoluble in carbon disulphide, and not poisonous. It ignites at about 260 C. in air. Red phosphorus unites with the halogens only when heated. Red phosphorus is prepared by heating the ordinary form in closed iron tubes to 300 C. A small amount of the yellow phosphorus remains unchanged ; this is removed by means of carbon disulphide, in which the red variety is insoluble. When a given amount of red phosphorus is burned, there is much less heat liberated than with an equal amount of the yellow form ; the red has therefore much less energy than the yellow. This statement agrees with the known fact that when yellow phosphorus is changed into the red there is an evolution of heat. MATCHES. 281 . 295. Molecular Mass of Phosphorus. The weight of a liter of phosphorus in the vapor condition is almost four times that of a liter of oxygen at the same temper- ature and pressure ; consequently the molecular mass of phosphorus, as a vapor, must be about 124, that is, about four times the atomic mass. This can be due only to the fact that the molecule of phosphorus in the gaseous condition contains four atoms. The molecular formula of phosphorus vapor is thus written P 4 , just as that of oxygen is O 2 . 296. Matches. Most of the phosphorus that is made is used to tip matches. The ordinary friction match, as made at present, consists of a splint of wood tipped, first, with sulphur, and then with a mixture con- taining some oxidizing agent, phosphorus, and an ad- hesive substance, like glue. The oxidizing agent may be potassium nitrate or chlorate, or the oxide of lead known as red lead, which has the formula Pb 3 O 4 . The chemical operations involved in lighting a match are essentially as follows : (1) The heat generated by rubbing the tip of the match against a rough surface causes the phosphorus to combine with the oxygen of the oxidizing agent in immediate contact with it. (2) The combustion of the phosphorus causes the sulphur to be raised to the kindling temperature of sulphur. (3) The burning of the sulphur raises the temperature of the wood to the kindling point ; and the match burns. Safety matches have not the property of being easily 282 THE NITROGEN FAMILY. ignited when rubbed ; they require contact with a spe.- cially prepared surface. This surface is usually on the side of the match box, and consists of red phosphorus mixed with sand. The tip of the safety match generally contains antimony trisulphide (Sb 2 S 3 ), an oxidizing agent, and glue. 297. Hydrogen Phosphide (PH 3 ). Hydrogen phos- phide, orphosphine, is a colorless gas which, as ordinarily made, is spontaneously combustible. The common method of preparing it is to heat a mixture of yellow phos- phorus and a strong solution of sodium hydroxide. The equation is, 3 NaH 2 P0 2 + PH 8 . sodium hypo- phosphite. Hydrogen FIG. 58. The apparatus (Fig. 58) consists of a generating flask con- taining the phosphorus and the sodium hydroxide solution. The stopper of the flask has two holes, one for a tube from a hy- drogen generator and the other for a delivery tube ending SALTS. 283 under water. The air of the apparatus is first washed out by means of hydrogen (or illuminating gas) ; the gas is then cut off and the flask is heated. The escaping phosphine may be collected in a receiver, as shown in the figure, and this exposed to the air, or the bubbles of the gas may be allowed to escape through the water directly into the air. The material of the white smoke formed when phosphine burns is phos- phorus pentoxide, water, and phosphoric acid. The equation for the combustion of phosphine is, 2 PH 3 + 4 2 > P 2 5 + 3 H 2 0. Pure phosphine is not actually ignited until its tem- perature reaches 100 C. As the gas is ordinarily prepared, however, it contains small amounts of the vapor of another phosphide of hydrogen (P 2 H 4 ), which is spontaneously combustible, and which, therefore, ig- nites the phosphine. 298. Phosphonium Salts. Phosphine may be re- garded as ammonia, NH 3 , with its nitrogen replaced by phosphorus. Although similar to ammonia in composi- tion, phosphine is much less basic. The aqueous solution of phosphine is not alkaline at all ; the compound PH 4 OH can hardly be present in the solution. Phosphine can, however, be made to unite with the halogen acids. The compounds thus formed correspond with the ammonium salts of the halogens ; hence they are called phosphonium salts. Phospkonium bromide (PH 4 Br) and phosphonium iodide (PH 4 I) are much like ammonium bromide and iodide, respectively. 284 THE NITROGEN FAMILY. Phosphonium iodide is decomposed by soluble hydroxides, much as ammonium chloride is (c/. 142). This fact is evi- dent from the equations, NH 4 C1 + KOH - KC1 -f NH 4 OII (i. e., NH 3 -f II 2 O) ; PH 4 I + KOH - KI + Pir 3 -J- H 2 O. 299. Phosphides. The phosphides of the metals may be considered derivatives of hydrogen pjiosphide, just as sulphides are of hydrogen sulphide. The formula of calcium phosphide is Ca 3 P 2 ; of silver phosphide, Ag g P. Calcium phosphide is a white solid ; when it is treated with water or with hydrochloric acid, it gives off hydrogen phos- phide. The equation resembles that for the action of hydro- chloric acid upon ferrous sulphide. Both equations are given. Ca 3 P 2 + G HC1 - 3 CaCl 2 + 2 PIT 8 . FeS + 2 HC1 - Fed, + H a S. 300. Oxides of Phosphorus. Two common oxides of phosphorus are known, viz., the trioxide (P 2 O g ) and the pentoxide (P 2 O 5 ). Both are white solids. The weight of a given volume of phosphorus trioxide in the gaseous state is known to be twice that demanded by the formula P 2 O 3 ; consequently it is better to write the formula P 4 O 6 . Phosphorus pentoxide is formed when phosphorus burns in air or oxygen free from moisture. ; 4P + 50 2 =2P 2 6 . It is a bulky, white solid which has great attraction for HYPOPHOSPHOROUS ACID. 285 moisture ; when put into water it hisses like hot iron. The product is metaphosphorie acid, HPO 3 . Phosphorus pentoxide has been referred to as capable of decomposing anhydrous nitric acid and producing nitrogen pentoxide (cf. 166). It is the anhydride of metaphosphorie acid, as nitrogen pentoxide is of nitric acid. 301. Oxygen Acids of Phosphorus. Several com- pounds of phosphorus with oxygen and hydrogen are known. Three of these form a series like the oxygen acids of chlorine (e/. 106 and 278); they are, - Hypophosphorous acid, H g PO 2 ; Phosphorous acid, H 3 PO 3 ; Phosphoric acid, H 3 PO 4 . 302. Hypophosphorous Acid. Attention has al- ready been called to the fact that when phosphorus acts upon sodium hydroxide (cf. 297) it produces sodium hypophosphite (NaH 2 PO 2 ) as well as phos- phine. With barium hydroxide, barium hypophosphite, Ba(H 2 PO 2 ) 2 , is produced. The hypophosphites are salts of hypophosphorous acid. This is a monobasic acid (cf. 105). Its graphic H\ ^0 formula is P^. , only the hydrogen atom at- ^ tached to oxygen being ordinarily replaceable by metals. 286 THE NITROGEN FAMILY. HypophosphorouB acid is a powerful reducing ayent, owing to the ease with which it goes over into phosphoric acid. 303. Phosphorous Acid. Phosphorous acid is an intermediate product in the oxidation of hypophos- phorous acid. It is itself a reducing agent, owing to its ready oxidation to phosphoric acid. Its anhydride is phosphorus trioxide, P 2 O 3 . P 2 3 + 3 H 2 = 2 H 3 P0 3 . Phosphorous acid may be prepared by treating phos- phorus trichloride or tribromide with water (<;f. 271). PC1 8 + 3 H 2 O - H 8 PO 8 + 3 11C1. Ordinary phosphorous acid is dibasic / its graphic formula .0 is, therefore, II P OH. 304. The Phosphoric Acids. The three phosphoric acids are, (1) orthophosphoric acid, or simply phosphoric acid (H 3 PO 4 ); (2) pyrophosphoric acid (H 4 P 2 O 7 ), and (3) metaphosphoric acid (HPO 3 ). There is still another phosphoric acid, from which all of the three named may be supposed to be derived, />// loss of water ; this is normal phosphoric acid, P(OH) r/ It corresponds to normal nitric acid, N(OH) 5 . But while the ordinary form of nitric acid is HNO 3 , i. e., the molecule of the normal acid minus two molecules of Water, the phosphoric acid from which the phosphates PllEPAliATlOy OF THE PUOSPUOIHC AC ID 8. 287 are chiefly derived is the orthophosphoric acid, H 3 PO 4 . The molecule of this acid is the molecule of the normal acid minus only one molecule of water. P(OH).-H 2 = H 8 P0 4 . If the molecule of orthophosphoric acid loses a mole- cule of water, we have metaphosphoric acid, HPO 3 . Tins acid corresponds to nitric acid, HNO 3 . Pyrophosphoric acid is derived from orthophosphoric acid by loss of one molecule of water from two molecules of the orthophosphoric acid. 2H,P0 4 H,0 = H 4 P,0 T . The complete anhydride of all the phosphoric acids is phos- phorus pentoxide, P 2 O 6 . Orthophosphoric acid is tribasic, a fact expressed in ^,OH its graphic formula, O P OH. It therefore forms two acid salts (cf. 103) and a normal salt Thus, with sodium hydroxide we ma}' get,- (1) Sodium di-hydrogen phosphate, N"aH 2 PO 4 ; (2) Disodium hydrogen phosphate, N"a 2 HPO 4 ; (tt) Trisodium phosphate, Na 3 PO 4 . Salts like the first of these, in which only one-third of the hydrogen is replaced, are called primary phosphates ; those in which two-thirds of the hydrogen is replaced are called secondary phosphates. The normal salts are tertiary phosphates. 305. Preparation of the Phosphoric Acids. The best way to obtain orthophosphoric acid is to treat red 288 THE NITROGEN FAMILY. phosphorus with nitric acid, and to evaporate the result- ing solution. At the ordinary temperature the acid consists of colorless, deliquescent crystals. Pyropliosphoric add is best made by heating the ortho- acid for some time to 260 C. ; the meta- acid is made by heating the ortho- or the pyro- acid to 300 C. Metaphosphorie acid always results when phosphorus pentoxide dissolves in water. P 2 6 + H 2 = 2 HP0 3 . When the meta- acid is boiled with water it goes over into the ortho- acid. 306. Salts of the Phosphoric Acids. MetapJios- phates are obtained from primary orthophosphates only, by the loss of one molecule of water from every mole- cule of the orthophosphate. Thus, sodium metaphosphate, NaPO 3 , is obtained by heating sodium di-hydrogen phosphate. NaH 2 PO 4 = H 2 O -f NaPO 3 . The so-called " metaphosphate bead'''' is made by heating either sodium di-hydrogen phosphate or sodium ammonium hydrogen phosphate upon a loop of platinum wire. Sodium ammonium hydrogen phosphate (also called " microcosmic salt") has the formula NaNH 4 HPO 4 . When heated it first loses ammonia, like any ammonium salt of a "fixed" acid, giving sodium di-hydrogen phosphate. This then loses water. + H 2 O . USES OF THE PHOSPHATES. 289 When a secondary orthophosphate, e. g., Na 2 HPO 4 , loses water, one molecule of water must come from two molecules of the phosphate ; hence a pyrophosphate results. 2 Ka 2 HPO 4 H 2 O = Na 4 P 2 O 7 (sodium pyrophosphate). 307. Uses of the Phosphates. A knowledge of the relations between the phosphates is essential to an understanding of the reactions involved in making fertilizers and phosphorus. The phosphate found in bone-ash and in nature is normal calcium phosphate, Ca 3 (PO 4 ) 2 . This, however, is insoluble in water. To convert it into soluble form for the use of plants, the normal phosphate is treated with sulphuric acid. This changes it into primary calcium phosphate, Ca(H 2 PO 4 ) 2 , which is soluble. Ca 3 (P0 4 ) 2 + 2 H 2 S0 4 2. CaSO 4 + Ca(H 2 PO 4 ) 2 . Calcium sulphate is much less soluble in water than the primary calcium phosphate ; hence the two can be separated readily. The primary phosphate of calcium is used not only as a fertilizer, but also in making baking powders (cf. 206) and as a source of phosphorus. The old process of making phosphorus from a phosphate consists, (1) in changing the phosphate into metaphosphate, and (2) in reducing the metaphosphate. 290 THE NIT HOG EN FAMILY. The change of primary calcium phosphate, like that of the sodium salt, into metaphosphate, takes place on heating. Ca(H 2 P0 4 ) 2 = Ca(P0 3 ) 2 + 2 H 2 O. The reduction of calcium metaphosphate to phosphorus takes place when the metaphosphate is heated with charcoal or with charcoal and silica, SiO 2 (cf. 292). B. Arsenic. 308. Occurrence and Preparation of Arsenic. The element arsenic is found in nature both free and combined. Its chief ores are realgar and orpiment (As 2 S 2 and As 2 S 3 , respectively), the oxide (As 2 O 3 ), and arsenopyrite (FeAsS). Arsenopyrite is iron pyrites (FeS 2 ) with half of the sulphur replaced by arsenic. The element may be prepared by reducing the oxide with charcoal. As 2 O 3 + 3 C 2 As + 3 CO. 309. Properties of Arsenic. Arsenic forms com- pounds which correspond closely with the compounds of phosphorus. The element itself is, however, more metallic than phosphorus. It exists in at least two allo- tropic forms. The ordinary form of arsenic is gray, has a crystal- line structure, and is about 5.7 times as heavy as water. It is not at all malleable, but, on the contrary, crumbles to powder when struck. PROPERTIED OF ARSENIC. 291 When arsenic is heated to about 500 C. out of contact with the air, it sublimes, forming a yellow vapor. By comparing the weight of a known volume of this vapor with that of oxygen under the same conditions, it is found that the molecular mass of arsenic vapor is about 300. The atomic mass being 75, the molecule must contain four atoms ; hence the molecular formula is As 4 . Above 1,700 C., however, most of the molecules con- taining four atoms dissociate into simpler molecules of two atoms each, i. e., As 2 molecules. Arsenic begins to burn at about 180 C. to form arsenic trioxide, As 2 O 3 . The flame is bluish-white. Like phosphorus and antimony, arsenic unites with chlorine at the ordinary temperature to form the chloride, AsCl 3 . 2 As + 3 C1 2 = 2 AsCl 3 . The same substance is formed when arsenic trioxide, As 2 O 3 , is treated with concentrated hydrochloric acid solution. As 2 O 3 + 6 HC1 2 AsCl 3 + 3 H 2 O. Arsenic trichloride is a colorless liquid ; it is decom- posed by an excess of water, giving arsenious acid and hydrochloric acid. Cl HOH OH As Cl + HOH As OH + 3 HC1. X Cl HOH x OH Arsenic trichloride is thus like phosphorus trichloride, which with water gives phosphorous acid and hydro- chloric acid (ef. 303). ' 292 THE NITBOGEN FAMILY. 310. Hydrogen Arsenide. Arsenic combines with nascent hydrogen to form hydrogen arsenide or ursine, AsH 3 , a substance which corresponds with phosphine gas, PH g . Arsirie cannot, however, be made to unite with hydrobromic acid, etc., to give compounds resembling ammonium and phosphonium salts (^f. 298). The most common method of getting arsine (mixed with hydrogen) is to add an arsenic compound to a flask in which hydrogen is being generated; the nascent hydrogen unites with the arsenic of the compound. Marsh's Test. Advantage is taken of the properties of arsine to test for the presence of arsenic in any substance ; the test is known as Marsh's test. To a flask in which pure hydrogen is being generated (Fig. 59), there is attached a calcium chloride tube and a hard glass J FIG. 59. tube drawn out as shown in the figure. The hydrogen is allowed to pass off until the usual test ((/. p. 13) shows that all air has been removed. The jet of hydrogen is now lighted, AESENIC TRIOXIDE. 293 and, a few drops of the liquid to be tested for arsenic are added through the thistle-tube. If arsenic is present, the flame changes to a bluish-white color, and a cold piece of porcelain held in the flame receives a shiny, black deposit, called an " arsenic mirror." If the hard glass tube is heated, the arsine passing through it is decomposed, and an arsenic mirror appears in the tube. Here the arsenic may be identified by passing hydrogen sul- phide, H 2 S, through the heated tube. The same precautions must be taken to have all air removed as in the case of hydro- gen. Hydrogen sulphide changes the arsenic into arsenic trisulphide, As 2 S 3 ; this is a golden-yellow solid called orpiment (from auri pigmentum). If, now, dry hydrochloric acid gas is passed through the tube, the arsenic trisulphide does not change. These properties serve to distinguish between the arsenic mirror and that of antimony (cf. 318). 311. Arsenic Trioxide. The oxides of arsenic are the trioxide (As 2 O 3 ) and the pentoxide (As 2 O 5 ) ; these correspond to the phosphorus oxides. Arsenic trioxide (often called "arsenic" or "white arsenic") is the most common arsenic compound. It is a white powder which sublimes, without melting, at about 220 C. The vapor has a garlic odor. When the vapor solidifies, the arsenic trioxide appears in the form of a transparent mass. At very high temperatures the molecule of the vapor is represented by As 2 O 3 ; but between 220 and 700 C. the molecules are doubled, and the formula is As 4 O 6 . Uses. Arsenic trioxide is used in medicine and as a poison. Its poisonous action upon the human system is rather slow, owing to its dissolving only slowly in the 294 THE NITROGEN FAMILY. liquid of the stomach. The antidote is a mixture of ferric hydroxide [Fe(OH) 3 ] and magnesia (MgO). The people of certain mountain districts are addicted to the use of arsenic trioxide because it enables them to breathe more easily when climbing. By beginning with very small quanti- ties and gradually increasing the dose, they are able to take much more than the lethal dose without injury. But the diffi- culty comes when they try to leave off the habit ; for they then suffer all the effects of arsenic poisoning. 312. Arsenious Acid. Arsenic trioxide is slightly soluble in water; the solution probably contains arseni- ous acid, HoAsO q . o rf As. 2 O 3 + 3 H 2 O = 2 H 3 AsO 3 . Arsenious acid is not known in the free state because it breaks up into arsenic trioxide (its anhydride) and water. The salts of arsenious acid are called arsenites. Solutions of these are formed when arsenic trioxide is treated with alkalies. Thus, sodium hydroxide and arsenic trioxide (with water this is arsenious acid) form sodium arsenite, Na 3 As() 3 . H 3 AsO 3 + 3 NaOH - Na 3 AsO 3 + 3 H 2 O. Many arsenites are derived from nwtarsenious acid, HAsO 2 , which may be looked upon as arsenious acid minus water. IT 3 AsO 3 = TT As(X + II 2 O. Sodium metar senile would be Double Nature of Arsenious Acid. Arsenic tri- oxide (or arsenious acid) reacts not only with alkalies, ARSENIC ACID. 295 giving arsenites and water, but also with acids, giving arsenic salts and water. Thus, with concentrated hy- drochloric acid and arsenic trioxide, we get arsenic tri- chloride and water. As 2 O 3 + 6 HC1 2 AsCl 3 + 3 H 2 O ; or, H 3 AsO 3 + 3 HC1 > AsCl 3 + 3 H 2 O. Arsenious acid has thus a double nature ; for toward strong bases it acts like an acid, forming with the base an arsenite; while toward a strong acid it acts like a base, giving with the acid a salt and water. Arsenic is, in fact, intermediate between the non-metals and the metals. Arsenic Greens. At least two arsenic compounds have a bright green color ; these are c-oppe* arsenite, called " Scheele's green" and a mixture of copper arsenite and copper acetate, called " Schweinfurt's green." Both of these are sold as ki Paris green." These dyes were formerly used to color wall- paper, paints, etc. They are too dangerous, however, and are now rarely used as dyes. Paris green is used to destroy potato- bugs, etc. 313. Arsenic Acid. Arsenic pentoxide, As 2 O 6 , is the anhydride of arsenic acid, H 3 AsO 4 . As 2 5 + 3 H 2 = 2 H 3 As0 4 . Arsenic acid is formed by dissolving arsenic or arsenic trioxide in concentrated nitric acid. (Compare the preparation of phosphoric acid from phosphorus, 305.) 296 THE NITROGEN FAMILY. The arsenic acids have formulas of the same type as those of phosphorus : H 3 AsO 4 is orthoarsenic acid ; H 4 As 2 O 7 is pyroarsenic acid ; HAsO 3 is metarsenic acid. Metarsenic acid finally gives, by loss of water, arsenic pentoxide. 2 HAsO 3 = H 2 O + As 2 O 5 . The arsenates are like the corresponding phosphates, 314. Arsenic Trisulphide. When the solution of an arsenic compound is treated with hydrogen sulphide, a yellow precipitate is generally produced ; this consists of either the trisulphide (As 2 S 3 ) or the pentasulphido (As 2 S 5 ). Both sulphides react with ammonium sul- phide [(NH 4 ) 2 S] and other soluble sulphides. The so- lution contains a sulpharsenite. or sulphar senate. Thus with sodium sulphide and arsenic trisulphide the equa- tion is, 3 Na 2 S + As 2 S 3 = 2 Na 3 AsS 3 . The sulpharsenite (Na 3 AsS 3 ) is simply an arsenite with its oxygen replaced by sulphur. When the sulpharsenite is treated with a dilute acid, e. g.. hydrochloric acid, sulphar senious acid, H 3 AsS 3 , is set free ; this breaks up into hydrogen sulphide and arsenic trisulphide. Arsenic trisulphide, being insoluble, is reprecipitated. These facts are shown in the equations, PHYSICAL PROPERTIES. 297 + 3 HC1 - H 3 AsS, + 3 NaCl. 2 H 3 AsS 3 = 3 H 2 S -f As 2 S 3 . Ammonium sulpharsenite, (NTf 4 ) 3 AsS 3 , undergoes a similar decomposition. C. Antimony. 315. Preparation of Antimony. Antimony is found in nature chiefly combined with sulphur in the mineral stibnite, Sb 2 S 3 . To obtain the element the sulphide is first roasted, i. e.,. heated in a stream of air, and then heated with charcoal. Roasting converts the antimony sulphide into the tnoxide (Sb 2 O 3 ), or the tetroxide (Sb 2 O 4 ), and sulphur dioxide. 2 Sb 2 S 3 + 9 2 - 2 Sb 2 3 + G SO 2 ; or, Sb 2 S 3 -f 5 O 2 - - Sb 2 O 4 -f 3 SO 2 . The reduction of either of these oxides by charcoal gives anti- mony and carbon monoxide. Sb 2 O 3 + 30 - 2 Sb + 3 CO. 316. Physical Properties. Antimony is a solid having a bright, silvery luster which is not easily tar- nished in air. Antimony is not malleable. At about 430 C. it melts to a liquid of a slightly higher specific gravity. When melted antimony solidifies it expands again; hence antimony is valuable as a constituent of materials for casts, such as type-metal. The specific gravity of. the solid is 6.7. The specific gravity of the yapor shows that in the gaseous condition the formula of the molecule is Sb 2 ( to i i CO M CO GO i ' GO CO to (1 h- ^ P ' M p- cr ? P N H 03 R K M O to CO to to o i t * i 1 CO P 05 to GO to 00 1 I to o te * W ^d w- ^J 8 CO ^ ^ O P o^ cr ac ^ ^ V ^-\ ^ s W to o GO 1 J GO CO to o cp 2 CO 1 I 1 I b Cn w ' 3 cj Si H CO K CO Cfc g s ^ to CO CO 1 I 00 *> to o CO to CO to 1 ' 05 W ^ p' W g 9 *J ^ ^ H-i O to 05 CO 1 ^ GO O P en CO Cn en 1 I CO M O -a w P 2 2 N3 1 I ""g CO ?$ ^en <^o v 1 I L_J !-* HH M Si(OH) 4 + 4 HF1. Instead of being set free, however, the hydrofluoric acid unites with some of the silicon tetrafluoride, forming fluosilicic acid, H 2 8iFl 6 . The name " fluosilicic acid " means silicic acid, H 2 SiO 3 , with its oxygen replaced by fluorine, three bivalent oxygen atoms by six univalent fluorine atoms (c/. 314, sulph- arsenites, and 194, thio sulphates). Many fluosilicates are known; potassium fluosilicate, K 2 SiFl 6 , is one of the few diffi- cultly soluble potassium salts. 318 SILICON AND BORON. 338. Silicon Carbide. Silicon carbide (SiC) or car- borundum is among the three hardest substances known, the others being boron carbide and the diamond. It is made by heating a mixture of powdered quartz, coke, saw-dust, and common salt in the electric furnace. Carborundum is not attacked by acids nor by solutions of alkalies. It burns only with great difficulty. Because of its hardness, powdered carborundum is used as a polishing and cutting agent. 339. Silicon Dioxide. Silicon dioxide is found as sand, quartz, etc. (c/. 335). The pure substance may be prepared by heating silicic acids, H 4 SiO 4 , H 2 SiO 3 , etc. H 4 Si0 4 = Si0 2 + 2 H 2 0. H 2 Si0 3 = Si0 2 + H 2 0. Silicon dioxide is thus silicic anhydride. When any form of silicon dioxide is fused with sodium or potassium hydroxide or carbonate, sodium or potassium silicate results. 2 NaOH -f SiO 2 Na 2 SiO 3 -f H 2 O. Ka 2 CO 3 + SiO 2 Na 2 SiO 3 + CO 2 . Calcium carbonate acts in the same way. When silica is fused with the oxide of a metal, a silicate is also formed. CaO -f SiO 2 = CaSiO 3 . Acids do not act upon silica (except hydrofluoric acid as in 337). 340. Silicic Acid. When a soluble silicate is treated with hydrochloric acid, a bulky mass, like gelatine, is SILICATES. 319 precipitated. This probably consists of normal silicic acid, H 4 SiO 4 . When the gelatinous mass is dried at the ordinary temperature, it loses water and becomes a non-crystalline powder. This is probably ordinary silicic acid, H 2 SiO 3 . When the powder is heated to a high temperature it loses water, forming silica, SiO 2 . The equations are given in 339. Polysilicic Acids. Besides the normal and ordinary forms of silicic acid, many other forms are possible ; these are known as polysilicic acids. They are derived from normal silicic acid by the loss of different proportions of water. A general for- mula for them all would be, o;Si(OH) 4 -yH s O. Thus, if x = 2 and y = 1, \ve have, 2H 4 Si0 4 H 2 = H 6 Si s 7 - Among the varieties of amorphous silica found in nature are agate, chalcedony, opal, cornelian., flint, amethyst, etc. These all contain water, and may therefore be looked upon as forms of the polysilicic acids. The colors of these substances are usually due to traces of impurities. 341. Silicates. The silicates are salts of silicic acid. The mineral chrysolite is magnesium silicate, Mg 2 SiO 4 . Serpentine is Mg 3 Si 2 O 7 . These are salts of the acids H 4 SiO 4 and H 6 Si 2 O 7 , respectively. Potassium, sodium, and calcium silicates are derived from the acid H 2 SiO 3 . Potassium silicate is known as "water glass " ; it is ued to make cements and artificial stone. Kaolin is practically pure aluminum silicate, Al 2 (SiO 3 ) 3 . It fuses only at a very high 320 SILICON AND BORON. temperature. It is used for making china and crockery ware, fire-bricks, etc. Clay, which is impure aluminum silicate, melts lower than kaolin ; it is used for making pottery, bricks, etc. The red color of baked clay is due to ferric silicate, Fe 2 (SiO 3 ) 3 . 342. Glass. Glass is a mixture of certain silicates, generally of sodium or potassium silicate with calcium or lead silicate. The silicates of calcium, lead, etc., crystallize when they sol- idify ; a glass made from them would break into fragments on cooling. The silicates of sodium and potassium, however, not only do not crystallize themselves, but even prevent the other silicates from crystallizing. Ordinary, soft glass, such as is used for window panes and bottles, is essentially a mixture of the silicates of calcium and sodium; it may be made by melting to- gether silica, calcium carbonate, and sodium carbonate, in the proper proportions. Hard glass is a mixture of the silicates of calcium and potassium. It is used for making chemical appara- tus, lamp globes, etc. Flint glass, such as is used in making optical instru- ments, etc., is a mixture of potassium and lead silicates. Enameled or * * milky ' ' glass is made by adding cryolite (cf. 266) to ordinary glass. " Granite ironware " or "porcelain-lined" ware con- sists of iron covered with an easily fusible glass, called enamel. Color is imparted to glass by the addition of small amounts PREPARATION OF BORON. 321 of other substances. Thus, a cobalt compound colors glass blue ; a cuprous compound, red ; a chromic compound, green. The etching of designs on glass is done with hydrofluoric acid, as described in 268. In certain kinds of etching a blast of sand is used. Pressed glassware is made in molds ; in cut glassware the designs are ground or polished by means of emery, carborun- dum, or sandstone wheels. All articles of glass, to be permanent, must be an- nealed. Annealing consists in allowing the hot object to cool regularly, so that its molecules may assume per- manent positions with reference to one another. Unan- nealed glass may fly to pieces at the slightest jar. B. Boron. 343. Occurrence of Boron. The element boron is the first member of the aluminum group of elements ( B(OH) 3 ,-f NH 3 . With chlorine, boron forms boron trichloride, BC1 3 , a colorless liquid. This is decomposed by water, giving boric acid and hydrochloric acid (c/. 303 and 337). BC1 3 + 3 H 2 > B(OH) 3 + 3 346. Boric Acid. Boric acid is made by adding concentrated hydrochloric or nitric acid to a hot solution of borax. It is a wliite, crystalline solid. Its aqueous solution has a faintly acid reaction with litmus, but colors turmeric paper brown, just as sodium hydroxide does.. Boric acid is in some respects like aluminum liy- droxide, which is an acid or a base, according to circum- stances. EXERCISES. 323 Boric acid is volatile with steam, as was indicated in 343 t Its solution in ethyl alcohol burns with a bright green flame. AVhen boric acid is heated it loses water, giving finally the an- hydride, B 2 O 3 . This redissolves readily in water, giving the acid. 347. Borax. The ordinary borates are derived not from the normal acid, H 8 BO 8 , but from tetraboric acid, H 2 B 4 O 7 . This is related to the normal acid just as the poly silicic acids are to normal silicic acid ( Ag + N0 3 + Na + CIO,. Another illustration : Ferrous chloride, FeCl 2 , gives with ammonium sulphide, (NH 4 ) 2 S, a black precipitate of ferrous sul- phide, FeS ; but ammonium sulphide does not precipitate the iron of potassium ferrocyanide solution, K 4 Fe(GN") 6 , because iron ions are not present. K 4 Fe(CN) 6 - 4 K + Fe(CK) 6 . 353. Electrolysis. What we call electrolysis is not the tearing apart of molecules by the electric current, but the carrying of electric charges by the ions from one electrode to the other. When we electrolyze dilute sul- phuric acid, as in 37, we have, before the introduction of the electrodes into the acid, ionization of most of the sulphuric acid molecules. H 8 SO 4 - ^ 330 DISSOCIATION AND MASS ACTION'. The current passes because the hydrogen ions carry -J- charges from the -j~ to the electrode, and the sul- phuryl (SO 4 ) ions charges from the to the -f- elec- trode. The bivalent >S'0 4 ion can carry as great a charge as two univalent hydrogen ions. + The charged ions (H and SO 4 ) are not atoms ; neither have they all the properties of free molecules. When they reach their respective electrodes, however, they take on their ordinary properties. The hydrogen ions give up their .-f- charges at the electrode, and become neutral atoms. These unite to form hv- drogen molecules. The sulphuryl ion gives up its charge at the -j- electrode, and becomes the radical SO 4 . This, being unstable, acts upon water, giving sulphuric acid and atomic oxygen. SO 4 -f H 2 O II 2 SO 4 + O . The oxygen atoms unite to give molecules of oxygen (O 2 ) or of ozone (O 3 ; cf. 287). 354. Hydrolysis. While strong acids, like hydro- chloric acid, etc., are much dissociated in solution, weak acids, like carbonic acid, hydrocyanic acid, etc., are dis- sociated only slightly. Consequently, when we dissolve the salt formed from a weak acid and a strong base, e. g., sodium carbonate, we get the reaction of the base, i. e., of OH ions. That this must be so is seen from the equation, + - + + - 2 Xa + C0 3 + 2 II + 2 Oil II 2 CO 3 + 2 Na + 2 OH. The carbonic acid is only slightly dissociated ; hence relatively few hydrogen ions remain in the solution, while the hydroxyl MASS ACTION. 331 ions are present in excess. Therefore sodium carbonate solu- tion reacts alkaline. The dissociation of the salts formed by weak bases and strong acids is similar ; only in this case we get the reac- tion of the acid, i. e., of hydrogen ions. Thus, when ferric chloride is dissolved in water, the ionic equation is, Fe + 3 Cl + 3 H + 3 OH Fe(OH) 8 -f 3 II + 3 Cl. Here the ferric hydroxide, Fe(OH) 3 , being a weak base and little dissociated, removes many OH ions from the "sphere of ac- tion " ; hence the H ions determine the reaction of the solution. Ferric chloride solution reacts acid. Morever, its rusty color shows that much undissociated ferric hydroxide (iron rust) is present in it. The illustrations just given are cases of hydrolysis, i. e., of " decomposition by water." 355. Mass Action. The recombination of dissoci- ated particles, whether in gaseous form or in solution, is influenced by the frequency with which the particles meet. If, therefore, we wish to stop or to diminish the dissociation of a substance AB, we simply see to it that a large excess of one of the dissociated particles, e. g., A, is present. The active mass of A, that is, the mass of it that can combine with B, is thus increased. The effects due to an excess of dissociated particles are cases of Mass Action. To illustrate : When we attempt to get the molecular 332 DISSOCIATION AND MASS ACTION. mass of phosphorus pentacJdoride, PC1 6 , by vapor density methods, we find that the molecular mass is too low, owing to the dissociation of some of the molecules into phosphorus trichloride and chlorine. PC1 5 PC1 3 + C1 2 . The dissociation can be prevented if the vapor density determination is carried out in an atmosphere of phosphorus trichloride ; because the chlorine molecules then meet mole- cules of phosphorus trichloride so frequently that practically no chlorine molecules remain free. Similar effects occur in solution. Thus, if \ve wish to pre- cipitate the sulphuric acid, i. e., the S0 4 ions, contained in a solution, we use barium chloride solution. If we use exactly the theoretical amount of barium chloride, however, much ba- rium sulphate remains in solution (as J3a and SO 4 ions). The reaction is only partly represented by the equation, Ba + 2 Cl + 2 II + S0 4 BaSO 4 + 2 II + 2 Cl, for the reverse operation also goes on. i BaSO 4 Ba -f SO 4 . Whatever will cause the $0 4 ions to meet more Ba ions in a given time will increase the precipitation of the $0 4 ions as barium sulphate. We obtain this result by adding a large excess of barium chloride solution. 356. Exercises. i. What substances are present in a solution of potassium nitrate in water ? Why is the reaction neutral ? EXERCISES. 333 2. Make a definition of an acid in terms of the ionization theory. Of a salt. 3. The reaction of a solution of potassium cyanide, KCN, is alkaline ; so is that of a solution of disodium hydrogen phos- phate, Na 2 HPO 4 . Explain. 4. Write the ionic equation for the neutralization of hydro- bromic acid, HBr, by potassium hydroxide. Explain. 5. Cold, concentrated sulphuric acid does not act upon zinc, but the dilute acid does." Explain. 6. Explain, in terms of the ionic theory, the action of hydro- yen sulphide, H 2 8, upon cupric sulphate solution (c/. 182). 7. To precipitate all the manganese of a manganese sulphate solution as sulphide ( cf. 182), we use a decided excess of am- monium sulphide, (KH 4 ) 2 S. Why ? CHAPTER XXIII. METALS. 357. Metals and Non-Metals. In our previous work we have studied chiefly non-metallic elements. There is, however, 110 sharp distinction between metals and non-metals, but rather a gradual change from one class to the other (cf. 327). But just as oxygen, chlorine, and sulphur have a distinct character, which no one would mistake for that of a metal, so there are certain elements having a typical metallic nature. Metals are usually opaque, and their polished surfaces are good reflectors of light ; hence the metallic luster. They are good conductors of heat and electricity. With oxygen and hydrogen the metals form bases; and by replacing the hydrogen of acids, i. e., ionic hydrogen, they form salts. Some elements, e. g., antimony, are both acid-forming and base-forming ; they are often called metalloids. 358. Occurrence of Metals. --The solid elements and compounds found in nature are called minerals. The minerals and mixtures of minerals from which metals are obtained are called ores. Some metals, e. g., gold and copper, occur free, that is, uncombined with other elements ; but most metals are found as oxides or 334 PROPEBTIES OF THE METALS. 335 sulphides. Some metals are found as carbonates, hydrox- ides, etc. 359. Extraction of Metals from Their Ores. If metals occur free they may be separated by mechanical means from minerals mixed with them. An illustration is the crushing of an ore in a stamp-mill and the wash- ing away of the lighter materials. Copper and gold are extracted in this way, although chemical methods are employed with inferior ores of these metals. The most common method of extracting metals is to reduce the oxide with carbon (charcoal or coke). This is the case with iron. If the ores used are not oxides, they are usually converted into oxides by " roasting " (cf. 178). Sulphides, hydroxides, and carbonates may thus be changed into oxides. Another method of reducing an oxide is to heat it in a stream of hydrogen. The oxygen and hydrogen unite and escape as steam, while the metal is left. This is a good laboratory method; but it is too expensive for commercial use. Chlorides are sometimes reduced by heating them with sodium. Aluminum was formerly obtained in this wa} r from its chloride. Several of the metals, e. g., aluminum, are obtained by the action of the electric current upon some of their compounds. 360. Properties of the Metals. Besides the gen- eral properties already mentioned, the metals possess 336 METALS. other properties in varying degrees. Thus, some metals, e. g., sodium and lead, are soft ; while others, e. g., chrom- ium and manganese, are hard. Sodium and lithium are light enough to float on water, while gold is 19.3 and platinum 21.5 times as heavy as water. Again, some metals evolve much energy in uniting with oxygen, while others form very unstable oxides. Usually the lighter metals, such as sodium and potassium, are very active chemically, and form strong bases ; while the heavy metals , such as lead and gold, are much less active. CHAPTER XXIV. THE ALKALI METALS. 361. General Properties. The metals, like the non- metals, are generally studied in groups or natural fami- lies based upon similarity of properties. The Alkali group consists of the jive metals named below and -the radical ammonium, NH^ the compounds of which re- semble those of sodium arid potassium (cf. 148). These metals are called " alkali " metals because the two most important members of the group are contained in the alkalies, i. e., in sodium and potassium hydroxides. ELEMENT. SYMBOL. ATOMIC MASS. SPECIFIC GRAVITY. MELTING POINT. Lithium. Li 7 0.59 180 C. Sodium. Ka 23 0,97 95. 6 C. Potassium. K 39 0.87 62.5 C. Rubidium. Rb 85.4 1.52 38.5 C. Caesium. Cs 133 1.85 26. 5 C. 337 338 THE ALKALI METALS. All of these metals have a silvery white luster and are easily cut. In air they become coated with a layer of the oxide and the hydroxide ; if carbon dioxide is present, these pass into the corresponding carbonates. The alkali metals burn when heated in air, and decom- pose water at ordinary temperatures ; therefore none of them is found free in nature. The salts of these metals are practically all soluble in water. The properties of the alkali metals change in the order of the atomic masses, e. (/., the higher the atomic mass the lower the melting-point. The chemical activity and the electro-posi- tive character increase from lithium to ccesium. Caesium is the most electro-positive element known. 362. Lithium. Lithium is widely distributed in nature, but no mineral known contains a large propor- tion of it. It is found in minute quantities in most mineral waters, in many plants, and in the blood. Lithium is the lightest of the metals. Its salts color the Bunsen flame crimson. 363. Sodium.. Sodium occurs widely distributed and in large quantities, especially as sodium chloride, NaCl. This exists as. rock-salt and sea-salt, and in many mineral springs. Sodium silicate is found in many rocks ; the nitrate is Chili saltpeter. Large deposits of the sul- phate and carbonate exist. The ashes of plants grow- ing in or near the sea contain sodium carbonate, and were formerly the source of many sodium compounds. PREPARATION AND PROPERTIES OF SODIUM. 339 364. Preparation and Properties of Sodium. At present, sodium is prepared by the electrolysis of sodium hydroxide, or of sodium chloride. Strontium chloride or potassium chloride is added to the sodium chloride to lower the melting-point. Formerly the metal was made by heating a mixture of the carbonate and charcoal in the absence of air. Na 2 CO 3 -j- 2 C - 2 !Na -f 3 CO. The sodium which distilled off was first condensed to a liquid, and was then collected under petroleum. Sodium is a white, soft metal that can be moulded be- tween the fingers, and can be readily pressed into wire. Although so soft at ordinary temperatures, it is quite hard at 20 C. It decomposes water, producing sodium hydroxide and hydrogen (cf. 46) ; this method is used to prepare pure sodium hydroxide. Sodium unites readily with oxygen (cf. 289), producing a mixture of the monoxide (Na 2 O) and the peroxide (Na 2 O 2 ). The properties of the pure monoxide are not known. The per- oxide has recently come into general use as an oxidizing and bleaching agent and as a source of free oxygen. It is decom- posed by water, giving, finally, sodium hydroxide and oxygen. (1) Ka 2 O 2 + 2 H 2 - 2 NaOH + H 2 O 2 . (2) 2H 2 2 = 2 An alloy of sodium and mercury, called sodium amalgam, is a useful reducing agent ; it is simply diluted 340 THE ALKALI METALS. sodium. An amalgam is an alloy of which mercury is one constituent. 365. Sodium Hydroxide. Sodium hydroxide, or caustic soda, is formed when sodium or either of its oxides acts upon water. Two of the commercial methods of preparing the hydroxide are as follows : - (1) Boiling a solution of sodium carbonate with slaked lime (calcium hydroxide). Na 2 CO 3 -|- Ca(OH) 2 > 2 NaOH + CaCO s . The calcium carbonate, being insoluble, is precipi- tated. The sodium hydroxide solution is drawn off and evaporated. (2) Electrolysis of a concentrated solution of sodium chloride. 2 ^aCl -f 2 H 2 2 NaOII + 1I 2 + C1 2 . The sodium hydroxide and hydrogen collect at the electrode. Sodium hydroxide is a white, deliquescent solid, very soluble in water. Both the solid and its solution absorb carbon dioxide readily. The solution has a soapy feel- ing, and turns red litmus blue. 366. Soap. Sodium hydroxide is one of the strong- est and most useful bases. When fats are boiled with it they are saponified, i. e., converted into soap. The fats are chiefly glyceryl salts of organic acids. They are decom- posed by sodium hydroxide into glycerine and the organic acid. The sodium salt of the organic acid is soap. SODIUM CARBONATE. 341 C 1 -lie. C 8 H S ^0 CO c"lC + HONa O CO C 17 II 35 glyceryl stearate sodium hydroxide .Oil C 3 H 5 OH -f 3 NaO. CO. C 17 H 8B . X OH glycerine sodium stearate (soap) The sodium stearate formed is "salted out" of solution by adding sodium chloride. When soap is dissolved in water it is partly hydrolyzed ( cf. 354 ) into sodium hydroxide and the organic acid ; the sodium hydroxide is the cleansing agent. "When soap is put into hard water, an insoluble scum is formed ; this is the calcium salt of the organic acid. 367. Sodium Carbonate. Sodium carbonate, or soda, is one of the most important chemicals manufactured. The principal methods of making it are the Solvay, or Ammonia, Process and the Le Blanc Process. The second of these is the more interesting historically, be- cause it was devised by Le Blanc for the French gov- ernment during the Revolution, when the supply was cut off; but the ammonia process is so much cheaper that fully three-fourths of the soda used is now made in this way. Both processes begin with common salt. The Solvay Process consists essentially in treating sodium chloride with ammonium hydrogen carbonate, NH 4 HCO 3 . KaCl + NH 4 HC0 3 NH 4 C1+ NaHCO s . This reaction takes place because sodium hydrogen carbonate 342 THE ALKALI METALS. (sodium bicarbonate) is not very soluble in water and is, there- fore, readily precipitated when ammonium hydrogen carbonate is added to concentrated salt solution. The solution of ammonium hydrogen carbonate is formed by passing carbon dioxide under pressure into a saturated solution of ammonia H 2 CO 3 - > NH 4 HC0 3 + H 2 O. Gentle heating converts the bicarbonate into carbonate. 2 NaHCOg = NajCO 3 + H 2 O -f CO 2 . The carbon dioxide thus set free is used again ; and nearly all the ammonia is recovered by heating the brine, from which the bicarbonate has crystallized, with slaked lime (cf. 142). The Le Blanc Process consists of essentially three opera- tions : (1) The conversion of common salt into sodium sulphate (cf. 92) ; 2 NaCl + H 2 S0 4 - Na 2 S0 4 + 2 HC1. (2) The reduction of sodium sulphate to sodium sulphide; Na 2 S0 4 + 40 - Ka 2 S -f 4 CO. (3) The conversion of sodium sulphide into the carbonate; 2 S + CaC0 3 - > Ka 2 CO 3 + CaS. The second and third operations take place together, sodium sulphate being mixed with limestone and coal-dust and the mixture heated. The sodium carbonate cannot readily be sepa- rated from calcium sulphide because both are soluble. If lime- stone is present in excess, however, some of it is dissociated into quicklime (CaO) and carbon dioxide ; the quicklime and the calcium sulphide form an insoluble compound. Soda comes into the market as calcined soda, or soda- ash, containing no crystal-water, and as crystallized soda, SODIUM CHLORIDE. 343 or sal sodae, Na 2 CO 3 . 10 H 2 O. Soda is used in great quantities in the manufacture of glass (c/. 342) and of sodium hydroxide ( 365). 368. Sodium Bicarbonate. Sodium bicarbonate is prepared by treating the normal salt with carbonic acid ( of sea-water, and is found in large deposits in Gali- cia (Austria), Germany, England, the United States, etc. In some places salt is mined as rock-salt, while in others the mixture of salt and earth is treated with water, the resulting brine being pumped to the surface and then evaporated. At Manistee, Mich., the brine is con- centrated and the salt continuously separated in a " vacuum " boiler. 344 THE ALKALI METALS. Sodium chloride crystallizes in colorless, transparent which decrepitate (cf. 50) when heated. It is only a little more soluble in hot than in cold water (cf. 59). Salt is necessary to the life of man and other animals, the hydrochloric acid of the gastric juice being derived from it (cf. 90). It is used in enormous quantities, not only as food, but as the starting material in the preparation of most com- pounds of sodium and of chlorine. 371. Sodium Nitrate. Sodium nitrate, or Chile saltpeter, occurs in enormous quantities in the Atacama Desert, in Chile (jrf. 163). It is very deliquescent, and s,o cannot be used for gunpowder, etc. It is con- verted into potassium nitrate, which is not deliquescent (/. 365). 2 KC1 -f 2 H 2 O > 2 KOH + H 2 -f C1 2 . It may also be made from potassium carbonate and " milk of lime," i. e., calcium hydroxide. 346 THE ALKALI METALS. - K 2 CO 3 + Ca(OH) 2 2 KOH -f CaCO z . Potassium hydroxide is a white, deliquescent solid It is a powerful base. It absorbs carbon dioxide from the air, forming potassium carbonate. 375. Potassium Carbonate, or Potash. Much potas- sium carbonate is prepared by the Le Blanc and Solvay processes from potassium chloride (c/. 367). The crude substance is obtained from wood-ashes. Anhydrous potassium carbonate is a powerful dehydrating agent (c/. 53). Potash is used chiefly in making the hydrjxide and hard glass (c/. 342). 376. Potassium Nitrate. Potassium nitrate (called, also, saltpeter and nitre) has already been described ( 5 KC1 + KC1O 3 -f 3 H 2 O. A cheaper way is to pass the chlorine into hot " milk of lime " ; calcium chlorate, Ca(ClO 3 ) 2 , is formed. This with po- tassium chloride gives potassium chlorate and calcium chloride. Ca(ClO 8 ) 2 -f 2 KC1 2 KC10 9 + CaCl 2 . POTASSIUM BROMIDE AND POTASSIUM IODIDE. 347 Potassium chlorate, being much less soluble in cold water than the other substances, crystallizes out, leaving the others in solution. Like the nitrate, potassium chlorate is valuable chiefly as an oxidizing agent. It is used in preparing oxygen, explosive mixtures, e. g., smokeless powder, matches, and fireworks. It is sold by druggists as " potash " for sore throats. * 378. Potassium Bromide (KBr) and Potassium Iodide (KI). Potassium bromide and potassium iodide may be prepared by the action of bromine and iodine, re- spectively, upon potassium hydroxide. The equations are analogous to the one for the action of chlorine upon caustic potash. 6 KOH -f 6 Br - 5 KBr + KBrO 3 + 3 H 2 O. 6 KOH -f 6 I - 5 KT -f KTO 3 + 3 H 2 O. By evaporating the solution containing bromide and bromate, or iodide and iodate, to dryness, and then heat- ing the residue sufficiently, we can decompose the brom- ate and the iodate just as we can the chlorate (cf. 19). 2 KBrO 3 = 2 KBr -f 3 O 2 . Potassium bromide and iodide are made, also, by treating the bromide and the iodide of iron with potassium carbonate. Fe 3 Br 8 -f 4 K 2 CO 3 - > Fe 3 O 4 -f- 8 KBr -j- 4 CO 2 . Fe 3 I 8 % 4 K 2 C0 3 - > Fe 3 4 -f 8 KI -f 4 CO 2 . 348 THE ALKALI METALti. The iron compounds are formed by adding bromine and iodine, respectively, to moist iron turnings. 379. Ammonium. - The formation of ammonium salts by neutralizing acids Avith ammonium hydroxide has already been described (cf. 148). The radical NH has not been isolated because it decomposes into ammonia and hydrogen. One of the many arguments for its existence is the formation of ammonium amalgam by the action of a strong solution of ammonium chloride upon sodium amalgam (cf. 364). Ka, Hg -f NTI 4 C1 !N"H 4 , Hg -f NaCl. sodium ammonium amalgam amalgam Ammonium amalgam is a bulky, metallic mass re- sembling sodium amalgam. It decomposes readily into ammonia, hydrogen, and mercury. 380. Exercises. 1. Why does the electrolysis of an aqueous solution of sodium chloride give sodium hydroxide and hydrogen at the - electrode (cf. 365) ? 2. Write the equations showing how potassium carbonate is made by the Solvay process (cf. 375). 3. All four substances given in the equation in 372 are soluble. Explain under what conditions the reaction can take place. 4. Describe the preparation of ammonium chloride, nitrate, and sulphate, giving equations. How does ammonium nitrate behave when heated ? Ammonium chloride ? EXERCISES. 349 5. What is the source of the ammonia of commerce? 6. How do you explain the fact that sodium bicarbonate solution reacts alkaline? 7. Which of the following gases would you dry with solid caustic potash : ammonia, hydrogen sulphide, carbon monoxide, carbon dioxide, oxygen? CHAPTER XXV. THE ALKALINE-EARTH METALS. 381. The Group. The "alkaline-earth" metals are so called because they form the transition from the alkalies to the " earth " metals, such as aluminum. In this chapter we shall consider glucinum (or beryllium), magnesium, calcium, strontium, and barium. The most important is calcium. 382. Calcium (Atomic Mass, 40) Calcium does not occur free ; but its compounds are found in large quan- tities. The most abundant is the carbonate, CaCO 3 ; this occurs as limestone, marble, chalk, calc-spar, and coral. The sulphate, CaSO 4 , the phosphate, Ca 3 (PO 4 ) 2 , and the fluoride, CaFl 2 , are also important miners. The metal is obtained by heating calcium Made with carbon in a stream of Irydrogen in an electric furnace. It is a silvery solid. It decomposes water like sodium and potassium. 383. Calcium Oxide (CaO). Calcium oxide is familiar to all as lime, or quicklime. It is made by heat- ing calcium carbonate in large furnaces called lime-kilns. CaCO 3 = CaO -f CO 2 . Lime is a white, amorphous solid, fusible only at the tem- perature of the electric furnace. For its use in the Drum- 350 CALCIUM HYDROXIDE. mond, or lime, light see 11. It absorbs water and carbon dioxide from the air, forming air-slaked lime, which consists Of the carbonate- and hydroxide. Lime is thus a good agent for removing water and carbon dioxide from gases (cf. 223). When lime is treated with a suitable quantity of water, the two unite to form a soft, dry powder. The operation is called " slaking " ; and the product, calcium hydroxide, is called " slaked lime." CaO + H 2 = Ca(OH) 2 . So much heat is evolved when water and lime unite, that lime improperly protected from water has been the cause of many fires. 384. Calcium Hydroxide Ca(OH) 2 . Calcium hy- droxide is not very soluble, less than 1% parts dis- solving in 1,000 of water. The solution is lime-water; its uses have been considered in 99 and 209. If more of the hydroxide is present then the water will dissolve, the liquid appears milky (" milk of lime " ; cf. 374 and 377). Uses. Calcium hydroxide is generally prepared from the oxide just before using. It is used to prepare ammonia (cf. 142), the hydroxides of sodium and potassium (cf. 365 and 374), bleaching powder (cf. 88 and 279), chlorates (cf. 377), glass, mortar, and cements (cf. 389); to purify suqar and illumi- nating gas (cf. 223) : to remove hair from hides ; to extract metals from their ores : as a disinfectant and white-wash, and in making stearin candles. 352 THE ALKALINE-EAETH METALS. 385. Calcium Chloride (CaCl 2 ) . Calcium chloride is usually made from calcium carbonate and hydrochloric acid (cf. 203). The anhydrous substance (made by heating at 200 C.) is very deliquescent, and dissolves in water with evolution of heat. It is used as a drying agent ( CaC0 9 Hence the " softening" of hard water by means of soda. 388. Calcium Phosphate, Ca 8 (P0 4 ) 2 . Normal cal- cium phosphate occurs as phosphorite; combined with calcium chloride or fluoride it forms apatite. Important deposits of these minerals are found in Florida and South Carolina. Calcium phosphate is the chief inor- ganic constituent of bones (cf. 292). Being insoluble, normal calcium phosphate is converted into the primary 354 THE ALKALINE-EARTH METALS. phosphate, Ca(H 2 PO 4 ) 2 , for use as a fertilizer ((/. 307). The primary phosphate is commonly known as " soluble phosphate " ; the mixture of the primary phosphate and calcium sulphate is called "super- phosphate." Phosphates Necessary for Plants. To be fertile, soil must contain calcium phosphate, an essential plant food. When the crops are removed, part of the phosphate of that region goes with them ; hence, phosphate must be returned to the soil if the land is to yield good harvests. If the crops are used as food for animals, part of the phosphate returns to the soil in manure ; if not, other fertilizers must be used. Nature usually keeps a soil fertile by means of decaying vegetation, which forms with the soil " vegetable mould." Fertilizers. A complete fertilizer supplies potassium, nitro- gen, and phosphorus. Most fertilizers, however, contain only one or two of these essentials. Potassium is usually returned to the soil as the sulphate or carbonate (wood-ashes ; cf. 378) ; sometimes as chloride. Nitrogen is frequently Supplied as ammonium salts, or as nitrates, especially sodium nitrate (cf. 371). Nitrogen is also contained in guano. Phosphorus is contained in fertilizers chiefly as "soluble phosphate," which is obtained by treating phosphate rocks or bone-ash with sulphuric acid. The" dry residue left after the waste products of slaughter- houses, e. silver ("coin silver"), and are said to be 900 fine. The grade 925 fine is called " sterling silver " ; British silver coins are of this grade. Silver-plating is usually done by electrolysis of the double cyanide of silver and potassium. A bar of silver forms the anode, and the object to be plated, the kathode ((/. 411). The PHOTOGEAPHT. 373 rough or "matt" surface is given the usual lustrous finish by polishing with chalk. The double cyanide solution is made by adding to a silver nitrate solution one of potassium cyanide (cf. 215) until the silver cyanide first precipitated is redissolved. Mirrors are made by precipitating silver upon glass. The silver is deposited from silver nitrate solution containing ammonia. This solution and a suitable reducing agent, e. (/., ammonium tartrate, or acetalclehyde (CH 3 CHO), are put upon the glass, and the glass is gently heated. The bright deposit of silver is washed, dried, and covered with varnish to protect it from the hydrogen sulphide, etc., of the air. 416.- Compounds of Silver. Silver nitrate, AgNO 3 , is the most important compound of silver. It is a white, crystalline solid, made by dissolving silver in dilute nitric acid. Silver nitrate is sometimes called lunar caustic. /Silver chloride, AgCl, silver bromide, AgBr, and silver iodide, Agl (cf. 285), are made by adding solutions of chlorides, bromides, and iodides, respectively, to solu- tions of silver salts. They are affected by light, and are used in photography. 417. Photography. Silver salts are used in photo- graphy because they change color and become insoluble in certain chemicals after being exposed to light. When a photographic plate is exposed in a camera, no change is visible until the plate has been developed. Developing consists in treating the plate with a reducing agent, such as ferrous sulphate, pyroyallic acid, hydroquinone, eikonogen, etc. When the plate is covered with the developing solu- tion, an image appears ; this is due to the precipitation of a 374 COPPER, SILVER, AND GOLD. film of silver, which produces variations of light and shade. Where the light acted strongly upon the plate, the deposit of silver is relatively heavy ; and where there was little action, there is little metal deposited. Just what the action of light upon the silver bromide of a plate is, is not definitely known, but it certainly makes the re- duction to silver take place more easily than is the case with ordinary silver bromide. The action of a developer may be illustrated by the following equation : 2 AgBr + C 6 H 4 (OH) 2 -f 2 KOH hydroquinone 2 Ag + 2 KBr -f- 2 H 2 O -f C 6 H 4 O 2 . quinone In this case the reduction of the silver bromide is due to the oxidation of hydroquinone to quinone. Fixing. When the plate is sufficiently developed, it is rinsed and put into a bath of sodium thio sulphate ( u hypo " ; cf. 194) to remove the silver salts not acted upon by light. This fixes the negative. Th^ plate is called a " negative " because in it dark objects appear light, and light objects dark. A " print " is made by placing the film of a sensitized paper next to the negative and exposing both so that the light passes through the negative. The image may appear on the paper at once ("printing-out" papers) or may have to be deirJojx'd ("developing" papers). In either case the prints are l *jLred" by removing the unchanged silver salt. Toning. Some papers are " toned " in a bath of gold, chloride, AuCl 3 , or platinum chloride, PtCl 4 , before fixing. Toning replaces part of the silver by gold or platinum. Blue prints are made on paper coated with a ferric salt (ferric ammonium citrate) and potassium ferricyanide, K 8 Fe(ON),.. After exposure, the picture is developed and fixed by washing METALLURGY OF GOLD. 375 it in waiter. The result is a blue print on a white ground. The process is used for copying plans, etc. 418. Gold (Atomic Mass, 197.2). Gold is found both native and combined. Even native gold is not pure, however, but contains silver, and often iron, copper, etc. The metal is frequently found enclosed in quartz or quartz-sand. Gold is obtained chiefly from Colorado and other western states, and from Australia, Siberia, and South Africa. The gold produced in the United States during 1899 was 3,437,210 fine ounces, worth $71,053,400. This was about one-fourth of the world's yield that year. 419. Metallurgy of Gold. Gold-mining is of two general kinds : (1) placer-mining, and (2) vein-mining. In placer-mining the clay and sand containing the gold are washed with water. The lighter particles are thus removed; while the gold and other heavy metals remain. Gold and silver are extracted from this mixture by mercury (r/V 413 ; amal- gamation^. Vein-mining consists in removing the gold-bearing rock from the earth and crushing it in stamp mills. The lighter materials are then washed away, and the gold is collected with mercury, as in placer-mining. Hydraulic-mining is a form of placer- mining done on a large scale with powerful streams of water. Instead of mercury, chlorine and bromine are used to re- move gold from the crushed ore. They form the soluble gold chloride, AuCl 8 , or bromide, AuBr s . This is extracted with water, and the gold is precipitated by means of charcoal or ferrous sulphate jftf. 420). The Cyanide Process depends upon the fact that gold is 376 COPPER, SILVER, AND GOLD. converted into the soluble double cyanide, KCN. AuCN, by a solution of the alkali cyanides. The gold is separated from the cyanide solution by electrolysis or by means of zinc. 420. Purification of Gold. The gold obtained by the processes described above is not pure. It can be separated from silver by adding aqua regia, which reacts with the gold. The solution is evaporated to remove nitric acid, the residue is dissolved in water, and the gold is precipitated by ferrous sulphate or some other reducing agent. 3 FeSO 4 + AuCl 3 Fe 2 (SO 4 ) 8 -f FeCl 8 -f Au. In the treatment of silver and gold with sulphuric acid ( alloy. Because of its malleability and weak chemical action, gold is much used by dentists for filling teeth. Gold-leaf is used in or- namentation. CHAPTER XXVIII. ALUMINUM (Atomic Mass, 27). 422. Occurrence of Aluminum. Although alumi- num does not occur free, it is the most abundant and widely distributed metal. Only oxygen and silicon are more abundant. Some of the most important minerals containing aluminum are feldspar (KAlSi 3 O 8 ), mica (KAlSiO 4 ), and cryolite (cf. 266). Granite is a mix- ture of quartz, feldspar, and mica. Clay results when granite and similar minerals are decomposed. All the other elements of the aluminum group are rare (c/. Periodic Table, 332). 423. Preparation. Aluminum was formerly pro- duced from anhydrous aluminum chloride and sodium.* A1C1 8 + 3 Na Al -f 3 NaCl. This method has been superseded by electrolytic processes, of which that of Hall (1887) is perhaps the most important. Hall's Process. The furnace used in the Hall process is a box of boiler iron (Fig. 61), the bottom and sides of which are lined with a mixture of coke and tar, rammed hard. The bot- tom forms the electrode, while the + electrode consists of forty large carbons suspended by copper rods. To begin the process, the carbons are lowered almost to the 878 PROPERTIES. 379 bottom of the furnace, cryolite is put in, and the current is turned on. The resistance to the passage of the current pro- duces heat enough to melt the cryolite. Pure, dry aluminum FIG. 61. oxide, A1 2 O 3 , is then mixed with the fused cryolite, and the electrolysis begins. The process is continuous, for the cryolite bath remains un- changed. The aluminum collects at the bottom, and is drawn off ; the oxygen of the aluminum oxide unites with the carbons to form carbon monoxide, which escapes. One company getting its power from Niagara produced 3,190 tons of aluminum dur- ing 1900 ; this was about one-third of the world's output. The aluminum oxide used is obtained from beauxite (or bauxite), A1(OH) 3 . The natural mineral has impurities, e. g., iron, silicon, etc. ; these are removed at present by fusing the beauxite with a little metallic aluminum. 424. Properties. Aluminum is a white, lustrous metal. Its specific gravity (2.6) is very low compared with that of other common metals, zinc being 7.1 and 380 ALUMINUM. iron 7.8 times as heavy as water. It melts at about 660 C. and vaporizes at the temperature of the electric furnace. Aluminum is a good conductor, is ductile and malleable, and has the tensile strength of cast- iron. Commercial aluminum is 95^ to 99. 6/0 pure. At white heat aluminum burns to aluminum oxide, A1 2 O 3 . Hydrochloric acid readily reacts with the metal, forming aluminum chloride. 2 Al + 6 HC1 - 2 AlCl 3 + 3 H a . Nitric acid and dilute sulphuric acid do not act upon it ordinarily. Aluminum reacts with solutions of salt and other chlorides if a little free acid is present. It reacts, also, with the hydroxides of sodium and potassium, form- ing aluminates and hydrogen. 6 NaOH + 2 Al - > 2 ]STa 3 AlO 3 -f 3 H 2 . Aluminum unites directly with the halogens, and with carbon, silicon, nitrogen, etc. 425. Uses. It is probable that more aluminum is used in the manufacture of iron and steel than for any other one purpose. The aluminum removes any oxygen the iron may have in combination, and thus increases the fluidity of cast-iron and steel. The next important use is as a conductor of elec- tricity. The Northwestern Elevated Railroad of Chi- cago is using twenty miles of 1^-inch aluminum cables, ALUMINUM OXIDE AND HYDROXIDE. 381 weighing 150,000 pounds, to transmit motive power to its trolley cars. Aluminum powder is used to reduce the oxides of many metals, e. g., chromic oxide, Cr 2 O 3 , and for flash-lights. A large part of the aluminum produced is used for kitchen utensils, scientific instruments, etc. ; the aluminum alloys also require much of the metal. The alloys with copper were described in 409. A new alloy, called magnalium, contains 75% to 90% aluminum and the remainder magnesium. 426. Aluminum Oxide and Hydroxide. Alumi- num oxide, A1 2 O 3 , occurs in the form of ruby, sapphire, and corundum. Emery, an impure form of corundum, is very hard, and is used for grinding and polishing. Aluminum oxide may be made by heating the hydroxide. 2 A1(OH) 8 = A1 2 O 3 + 3 H 2 O. Aluminum hydroxide, A1(OH) 3 , may be precipitated by adding ammonium hydroxide to the solution of an aluminum salt, e. g., aluminum chloride. Aluminum hydroxide reacts with both acids and alkalies (except ammonium hydroxide) ; with acids it gives aluminum salts, and with alkalies, aluminates (cf. 424). It is, therefore, a base toward strong acids and an acid toward strong bases (cf. 312 and 317). (1) A1(OH) 8 -f 3 HC1 > A1C1 8 -f 3 H 2 O. (2) H 3 A1O 3 -f '* NaOH Ka 3 AlO 3 + 3 H 2 O. Aluminum hydroxide unites either chemically or mechani- cally with many dye-stuffs, and also with certain fabrics. Ad- ALUMINUM. vantage is taken of this fact to " fix " the dyes in fabrics that do not readily hold color. The hydroxide is therefore called a mordant, from Latin to bite, because it bites into the fabric. 427. Aluminum Salts. Aluminum chloride, A1C1 3 , is soluble in water, and crystallizes with crystal- water (AlClg. 6 H 2 O) ; but it is so much hydrolyzed that when the water is expelled hydrochloric acid es- capes, and aluminum hydroxide remains. A1C1 8 + 3 H 2 - - A1(OH) 3 + 3 HC1. The anhydrous chloride is made by letting dry hydro- chloric acid gas act upon hot aluminum filings. It is a hygroscopic, white powder. Alums. Aluminum sulphate, A1 2 (SO 4 ) 3 , forms double salts with the sulphates of univalent metals, e. g., with potassium sulphate, K 2 SO 4 ; these double sulphates are called alums. Tho alums crystallize with twenty-four molecules of crystal-water. Potash-alum is K 2 SO 4 , A1 2 (SO 4 ) 3 , 24 H 2 O ; silver-alum is Ag 2 SO 4 , A1 2 (SO 4 ) 3 , 24 H 2 O. Other trivalent elements may replace aluminum. Thus N"a 2 SO 4 , Ee 2 (SO 4 ) 3 , 24 H 2 O would be sodium iron-alum. Chrome-alum is K 2 SO 4 , Cr 2 (SO 4 ) 3 , 24 H 2 O. Aluminum Carbonate and Sulphide. The elec- tro-positive properties of aluminum are so weak that even its salts with strong acids, e. g., the chloride and sulphate, are readily hydrolyzed. Its salts with weak acids cannot exist in the presence of water, being com- pletely decomposed into the hydroxide and the free acid. When the acid is volatile, it of course escapes. This is the case with the carbonate. When aluminum alts are treated PORCELAIN, STONEWARE, ETC. 383 with the solution of a carbonate, the products are aluminum hydroxide and carbonic acid. Hence carbon dioxide escapes. (1) 2 A1C1 3 + 3 N"aoCO 3 A1 (CO 3 ) 3 -f 6 NaCl. (2) A1 2 (C0 3 ) 3 + G H 2 6 2 A1(OII) 3 -f 3 H 2 CO 3 . With the sulphide the result is similar. (1) 2 A1C1 3 + 3 (NII 4 ) 2 S A1 2 S 3 + 6 NH 4 C1. (2) A1 2 S 3 -f 6 H 2 O 2 A1(OII) 8 -f 3 H 2 S. 428. Porcelain, Stoneware, Etc. Aluminum sili- cate, Al 2 (SiO 8 ) 3 , is essentially the substance from which porcelain, etc., are made. In a pure form it is kaolin / in an impure form, day (cf. 341). Porcelain is made by mixing white kaolin with more fu- sible substances, such as feldspar, shaping the plastic mixture into form, and heating it to a high temperature. The more fusible portion (the feldspar) melts, and cements the whole together. Porcelain is hard and translucent, and withstands the action of heat and chemicals better than glass, hence it is used for many purposes in chemical laboratories. Stoneware is opaque, for it has not been heated enough to make the feldspar penetrate the kaolin as much as in porcelain. }arthenware is made from common clay, hardened by heat, but not fused. It is glazed by putting common salt into the furnace at the time of heating. This forms a covering of sodium aluminum silicate over the porous sur- face. Bricks, tiling, jugs, terra- cotta, etc., are examples. Ultramarine is a blue coloring substance, made by melting together kaolin, sodium carbonate, and sulphur. This substance was once very valuable, but thousands of tons of it are now made every year. CHAPTER XXIX. IRON, COBALT, AND NICKEL. 429. Iron (Atomic Mass, 56). Iron is the most use- ful of the metals. It is also one of the most widely distributed, since it is found in many minerals, in the soil, and in natural waters. Iron is an essential part of chlorophyll, the green material of plants, and of the red coloring matter of the blood. It is present in meteor- ites, and in the sun and stars. The principal ores of iron are hcematite (Fe 2 O 3 ), magnetite (Fe 3 O 4 ), broivn iron ore [Fe 2 O 3 . 2 Fe(OH) 3 ], and siderite, or spathic iron, (FeCO 3 ). Iron pyrites, FeS 2 , is a source of sulphur. 430. Metallurgy. The ores of iron are reduced by heating them with carbon (coke or coal) in a blast-fur- nace (Fig. 62). A flux (limestone or feldspar) is added to combine with the ashes of the coal and form a slag ( 4 FeCl 8 + 2 H 2 O. If no acid is present, part of the iron is oxidized to ferric salt, and part to ferric hydroxide (rust). 12 FeCl 2 + 3 2 + 6 H 2 O > 8 FeCl 3 + 4 Fe(OH) 8 . Ferric chloride, FeCl 3 , is formed in solution by passing chlorine into ferrous chloride solution, or by treating iron with aqua regia, and evaporating repeatedly with hydro- chloric acid. The anhydrous salt is made by passing chlorine over red-hot iron. It looks like maple-sugar. 437. Iron Sulphates. A solution of ferrous sul- phate, FeSO 4 , results when iron reacts with dilute sul- phuric acid. The crystals known as green vitriol or " copperas " are FeSO 4 . 7 H 2 O. Green vitriol is used in making inks, in dyeing, and as a deodorizer. Fer- rous sulphate is oxidized like the chloride. With ammonium sulphate it forms the double salt, ferrous ammonium sulphate, (NH 4 ) 2 SO 4 . FeSO 4 . 6 H 9 O ; this is much less easily oxidized than ferrous sulphate alone. Ferric sulphate, Fe 2 (SO 4 ) 3 , is made, in solution, by oxidizing ferrous sulphate with nitric acid in the pres- ence of sulphuric acid. 438. Potassium Ferro- and Ferri-cyanides. Potas- shim ferrocyanide, K 4 Fe(CN) 6 , is called "yellow prussiate" COBALT. 391 (i. e., cyanide, cf. 216) "of potassium." It is a yellow, crystalline solid. With a ferric salt it produces Prussian blue. Potassium ferricyanide, K 3 Fe(CN) 6 , is obtained by oxidizing the ferrocyanide with chlorine. It is a red, crys- talline solid. Potassium /e?vocyaiiide may be written Fe(CN) 2 . 4 KON ; this formula shows that the iron is ferrous, i. e., bivalent. (Ferro- means ferrous.) In the /err/cyanide, Fe(CN) 3 . 3 KCN, iron is in the/erHc condition. 439. Nickel (Atomic Mass, 58.7). Nickel occurs with iron in meteorites. Its ores (silicates of nickel) are found chiefly in Canada, Norway, and New Caledonia. Like iron, it forms two classes of compounds ; these are nickelous and nickelic compounds. The common com- pounds are nickelous. Most of them are green. The formula of nickel sulphate is NiSO 4 ; that of the ni- trate, Ni(NO 3 ) 2 ; that of the sulphide, NiS. Nickel is used to plate other metals to protect them from the atmosphere. It is used, also, in making alloys, e. g., German silver, nickel steel, and the United States five-cent piece. Both nickel and cobalt are attracted by the magnet. 440. Cobalt (Atomic Mass, 59). Cobalt occurs com- bined with arsenic and sulphur, and often associated with nickel. Cobalt salts are red in solution or combined with much crystal-water, and blue when anhydrous or with little crystal- water. A solution of cobaltous chloride, CoCl 2 , is used as a sympathetic ink. CHAPTER XXX. MANGANESE AND CHROMIUM. 441. Manganese (Atomic Mass, 55). Manganese occurs chiefly as the black oxide, MnO 2 , the mineral pyrolusite. The pure metal is very hard, and fuses only at a high temperature. In some ways it resembles iron. Thus, it forms two series of salts, manganous and man- ganic salts, corresponding to ferrous and ferric salts. The manganous salts, however, are more stable than the ferrous salts, and are not readily oxidized. Manganic salts are much less stable than manganous salts. The latter are usually pink in color and crystalline. They are formed when the higher oxides of manganese are treated with acids, oxygen being either set free or else used in oxidizing the acid (ef. 81). 442. Manganese Oxides. Manganese forms the fol- lowing oxides : Manganous oxide, MnO ; Manganese dioxide, MnO 2 ; Manganic oxide, Mn 2 O 3 ; Manganese heptoxide, Mn 2 O 7 . Manganous-manganic oxide, Mn 3 O 4 ; All of these are solids except the last, which is a dark liquid. The oxide Mn 2 O 5 , corresponding to the mangan- ates, is not known. The most important oxide of manganese is the di- POTASSIUM PERMANGANATE. 393 oxide. This is used in preparing chlorine and oxygen, and in decolorizing glass. The manganese dioxide used in making chlorine is recovered by treating the manganous chloride produced (cf. 81) with slaked lime. This forms manganous hydroxide, Mn(OH) 2 . By means of steam, air, and more lime, this is converted into manganites having the formulas CaMnO 3 (i. e., CaO. MnO 2 ) and CaMn 2 O 5 (i. e., CaO. 2 MnO 2 ). Both of these give chlorine when treated with hydrochloric acid. 443. Potassium Permanganate. Manganese forms not only salts, in which the manganese is the electro- positive element, but also manganates and permangan- ates, in which the manganese has the same relation to the compound that sulphur has to the sulphates. The more highly oxidized the manganese is, the less basic does its oxide become. Manganese heptoxide, Mn 2 O 7 , is the anhydride of permanganic acid, HMnO 4 ; the most important salt of this acid is potassium permanganate, KMnO 4 . Potassium permanganate is formed by boiling the solution of the manganate, K 2 MnO 4 , and passing carbon dioxide or chlorine into it. (1) 3 K 2 Mn0 4 + 2 CO 2 = 2 K 2 CO 3 + 2 KMnO 4 + MnO 2 . (2) 2 K 2 MnO 4 + C1 2 = 2 KC1 + 2 KMnO 4 . The permanganate separates from solution in prisms. It col- ors water a deep purple. The manganate is obtained by fusing a mixture of manganese dioxide, potassium hydroxide, and an oxidizing agent, e. g., potassium nitrate or chlorate. A crude permanganate solution is used to oxidize sewage. 394 MANGANESE AND CHROMIUM. it is made by treating crude sodium manganate or potassium manganate with dilute sulphuric acid. 444. Oxidation by Permanganate. Most of the uses of potassium permanganate are due to its easy liberation of oxygen. There is a difference depending upon whether it acts in acid or in alkaline solution. We represent the oxidation of oxalic acid by potas- sium permanganate in the presence of sulphuric acid as" follows : - 2 KMnO 4 -f 5 H 2 C 2 O 4 + 3 H 2 SO 4 = 2 Mn$O 4 + K 2 $O 4 -f Another process depends upon the oxidation of a ferrous to a ferric compound. 2 KMn0 4 + 10 FeS0 4 + 8 H 2 SO 4 == 5 Fe 2 (SO 4 ) 3 + K 2 SO 4 + 2 MnSO 4 + 8 H 2 O. Botli of these reactions are sharp and complete, hence they have important uses in volumetric analysis. To understand the oxidizing action of potassium perman- ganate in acid solution we must look upon this substance as made up of potassium oxide and manganese heptoxide. 2 KMnO 4 = K 2 O + Mn,O 7 . The molecule of manganese heptoxide gives up five atoms of oxygen. Mn 2 O 7 = 2 MnO -f 5 O. The manganmts oxide then reacts with the acid to form a man- ganous salt and water, and the potassium oxide to form a po- tassium salt and water. CHROMIUM. 395 MnO + H 2 S0 4 = MnS0 4 + H 2 O. K 2 O + H 2 SO 4 = K 2 SO 4 -f HoO. The complete equation is, therefore, 2 KMnO 4 -f 3 H 2 SO 4 = 2 MnSO 4 -\- K 2 SO 4 + 3 H 2 O -f 5 O. When the permanganate is used in neutral or alkaline solution, the available oxygen is less than in acid solu- tion, e. y., - 2 KMnO 4 -f H 2 O = 3 O + 2 KOH -f 2 MnO 2 . If there is sufficient alkali, the manganese dioxide unites with it to form a manganite, e. g., K 2 MnO 3 , which re- mains in solution. The action of potassium permanganate with sulphurous acid and with hydrogen peroxide is represented thus : 2 KMn0 4 -f 5 II 2 8O 3 + 3 H 2 SO 4 = K 2 SO 4 + 2 MnSO 4 + 5H 2 SO 4 -f3H 2 O. 2 KMnO 4 + 5 H 2 O 2 + 3 H 2 SO 4 ="K 2 SO 4 + 2 MnSO 4 + 8 1I 2 O + 5 C) 2 . 445. Chromium (Atomic Mass, 52.1). Chromium is a comparatively rare element. It occurs chiefly as chromite, Fe(CrO 2 ) 2 or FeO. Cr 2 O 3 . The name of the element is from the Greek chroma, meaning color. Chromium compounds have many decided colors. The metal is prepared at a high temperature by re- ducing the oxide with carbon or the chloride with sodium. It is steel-gray, very hard, and very difficult to fuse. With iron it forms a hard alloy called "chrome- steel." 396 MANGANESE AND CHROMIUM. 446. Oxides and Hydroxides. The most important oxides of chromium are chromic oxide (Cr 2 O 3 ) and chromium trioxide (CrO 3 ). Chromous oxide, CrO, is the one from which the chromous salts are derived. Chromic oxide is a valuable pigment known as " chrome- green." It is made by driving off water from the hydroxide, Cr(OH) 3 , or by heating a mixture of potassium dichromate, ammonium qhloride, and sodium carbonate. Chromium trioxide separates as bright red crystals when strong sulphuric acid is added to a saturated solution of potas- sium dichromate. Although often called " chromic acid," it is really the an- hydride of chromic acid, H 2 CrO 4 . It is a powerful oxidizing agent. Chromic acid is unknown. Chromic hydroxide, Cr(OH) 3 , is a green solid, formed when an alkaline hydroxide, carbonate, or sulphide is added to a chromic salt solution. It is soluble in an excess of the alkali, forming a chromite (cf. 426). Chromites are derived from the substance having the formula HCrO 2 ; this is chromic hydrox- ide minus water. The carbonate and sulphide of chromium are decomposed by water, like the corresponding aluminum salts (cf. 427). 447. Chromous and Chromic Salts. Chromous salts are so readily oxidized that they are very hard to prepare. In this respect chromium differs decidedly from manganese, the corresponding salts of which, the manganous salts, are stable (cf. 441). The chief chromic salts are chromic chloride (CrCl 3 ) and chrome- alum, K 2 Cr 2 (SO 4 ) 4 . 24 H 2 O. Chromic chloride is obtained as a beautiful, lavender-colored DOUBLE NATURE OF CHROMIUM. 397 solid by passing chlorine over a mixture of chromic oxide and carbon. The carbon reduces the oxide, and, at the same time, ehlorine unites with the chromium. Chromic chloride is formed in solution by reducing a chrom- ate or dichromate with alcohol, hydrogen sulphide, sulphurous acid, etc. The hydrated salt is green. 2 K 2 CrO 4 + 3 H 2 S + 10 HC1 = 4 KC1 -f 2 CrCl 3 + 8 H 2 O + 3 S. Chrome-alum is a violet-colored, crystalline substance, formed as a by-product in certain operations in which potas- sium dichromate is used as an oxidizing agent. It is analogous to ordinary alum, but contains chromium instead of aluminum (c/- 427). 448. Double Nature of Chromium. Chromium is not only a metal, but also an acid-forming element. Its lower oxides, like those of manganese, form salts with acids, and are, therefore, basic ; but its higher oxides, especially the trioxide, CrO 3 , are the anhydrides of acids. The chromites formed by the reaction of chromic hydroxide with alkalies (ef. 446) are not important ; although ferrous chromite, Fe(CrO 2 ) 2 , is the chief chrom- ium ore. The ehr ornate* and the dickromateg, however, are the most important chromium compounds. Chrom- ium is in the same periodic group with sulphur ; chromic acid (H 2 CrO 4 ) corresponds with sulphuric acid and dichromic acid, H 2 Cr 2 O 7 or H 2 O. 2 CrQ 3 , with fuming or disulphuric acid, H 2 S 2 O T . 398 MANGANESE AND CHROMIUM. 449. Chromates and Bichromates. Potassium chrotn- ate, K 2 CrO 4 , is made by roasting ehromite, Fe(CrO 2 ) 2 or FeO. Cr 2 O 3 , with potassium , carbonate and quick- lime in the oxidizing flame of a reverberatory furnace (fjf. Fig. 60). On a small scale the ehromite is heated with a mixture of potassium nitrate and carbonate. To understand the formation of the chruiuate we must think of it as made up as follows : 2 K 2 O -f O 2 O g + 30=2 K 2 OO 4 . The potassium oxide is present in the carbonate, the chromic oxide conies from the ehromite, and the oxygen from the oxidiz- ing agent or the air. In general, we change a chromic com- pound to a chromate by oxidizing it in the presence of a base. Potassium chromate is yellow, like most chromates. Acids change it to potassium diehromate, K 2 O 2 O 7 . Potassium diehromate forms large, red crystals, which are soluble in about ten parts of water at the ordinary temperature. Alkalies change it to the chromate. To understand the relation between the chromates and dichromates we must look upon their molecules as made up as follows : K 2 Cr,0 7 = K 2 0. 2Cr0 3 . K 2 Cr0 4 = K 2 0. Cr0 3 . Addition of alkalies, e. gr., potassium hydroxide, to the di- ehromate produces chromate. K 2 0. 2 Cr0 3 + 2 KOH (f. e., K 2 O. H 2 O) = H 2 O -f 2 K 2 OO 4 (i. e., 2 K 2 O. CrO 3 ). With acids, however, the reverse change takes place. 2 K 2 O. CrO 3 -f H 2 SO 4 = K 2 O. 2 CrO 8 -f K 2 SO 4 -f H 2 O. OXIDATION BY CHBOMATES. 399 Sodium dichromate, Na 2 O 2 O 7 , is often employed in place of the potassium salt, owing to its greater solu- bility. Both the chromates and the dichromates are used in dyeing, in calico-printing, and as oxidizing agents. Po- tassium dichromate is used in photography. Solutions of chromates and dichromates precipitate many metals, e. g., lead, silver, and barium, as chromates. 450. Oxidation by Chromates and Dichromates. - When a chromate or a dichromate is used as an oxidiz- ing agent, the reactions are the reverse of those that take place when a chromate is synthesized : the chrom- ate (or dichromate) is reduced in the presence of an acid to a chromic salt. Thus we think of potassium chromate as breaking down into potassium oxide, chromic oxide, and oxygen. 2 K 2 CrO 4 = K 2 O + Cr 2 O 3 + 3 O. The oxygen is available for oxidation ; while the oxides unite with the acids, giving salts. If the acid present is sulphuric acid, potassium sulphate (K 2 SO 4 ) and chromic sulphate, Cr 2 (SO 4 ) 3 , are formed. These unite to produce chrome-alum, K 2 Cr 2 (SO 4 ) 4 . 24 H 2 O. If, however, hydrochloric acid is present, we get a mixture of the chlorides of chromium and potassium. If there is no other reducing agent present, the nascent oxygen attacks the hydrochloric acid, giving chlorine and water. K 2 Cr 2 7 + 14 HC1 = 2 KC1 + 2 CrCl 3 + 7 H 2 O -f 3 C1 2 . CHAPTER XXXI. LEAD, TIN, AND PLATINUM. A. Lead (Atomic Mass, 206.9.) 451. Occurrence and Preparation of Lead. Lead occurs chiefly as galena, or galenite, PbS ; and is ob- tained from it by the following process : The galena is first roasted in a reverberatory furnace (cf. Fig. 60). By this operation part of the ore is changed to the oxide, PbO, and part to the sulphate, PbSO 4 , while some remains unchanged. (1) 2 PbS -f 3 O 2 2 PbO -f 2 SO 2 . (2) PbS + 2 O 2 > PbSO 4 . After the oxidation has gone far enough, the furnace doors are closed, and the mixture is heated without the admission of more air. The lead oxide and sulphate then react with the un- changed sulphide as follows : (1) PbS -f 2 PbO 3 Pb -f SO 2 . (2) PbS0 4 + PbS 2 Pb -f 2 S0 2 . If the ores are poor, they are often heated with iron. PbS -f Fe = FeS -f Pb. If there is enough silver to pay for its extraction, the Parkes process is used (cf. 413). 452. Properties and Uses. Lead is soft, blue-gray metal having a high luster. It is malleable, but not very ductile. It melts at about 325 C. 400 COMPOUNDS OF LEAD. 401 Although lead is easily tarnished in the air, the cor- rosion does not penetrate, as with iron. Ordinary hard waters act but little upon lead ; but soft waters contain- ing carbon dioxide, organic matter, or much chloride or nitrate, attack it. Such waters should not be carried through lead pipes for household purposes. All compounds of lead are poisonous. Nitric acid acts readily upon lead, but hydrochloric and dilute sulphuric acids do not. Some metals, e. g., zinc, separate lead from the solutions of its salts. Lead pipes are used for conveying water and as sheaths for the cables of telephone wires. In sheet form the metal is used to line the " leaden chambers " (cf. 184), and the sides and floors of vats and tanks where certain chemical processes are carried on. Large quantities of^it are made into shot, bullets, type-metal, solder, pewter, and the plates of storage batteries. 453. Compounds of Lead. Several oxides of lead are known. Among these are the suboxide (Pb 2 O), the monoxide (PbO), the dioxide (PbO 2 ), and " red lead " or " minium," which is Pb 3 O 4 . Lead monoxide, or "litharge," is formed by heating lead in a current of air. At 400 C. it takes up more oxygen, forming " red lead," Pb 3 O 4 . When red lead is treated with dilute nitric acid, lead dioxide remains as a brown powder. Lead dioxide acts upon hydrochloric acid to give chlorine. Lead nitrate, Pb(NO 3 ) 2 , is made from litharge and nitric acid ; lead acetate, Pb(C 2 H 3 O 2 ) 2 , from litharge 402 LEAD, TIN, AND PLATINUM. and acetic acid. Both are white, crystalline solids. Lead acetate is called " sugar of lead." Lead sulphate (PbSO 4 ) and lead chromate (PbCrO 4 ) are insoluble in water. Lead chromate is called " chrome-yellow." Lead chloride, PbCl 2 , is difficultly soluble in cold water, but dissolves in hot water. Lead carbonate, PbCO 3 , is precipitated from solutions of lead salts by ammonium carbonate solution; the carbonates of sodium and potassium, however, give a basic carbonate instead. Basic lead carbonate is made by various methods, and on a large scale ; it is the pigment " white-lead." It forms a good paint, but turns brown or black in the presence of hydrogen sulphide. Lead sulphide, PbS, is precipitated from., the solution of a lead salt by soluble sulphides and by hydrogen sul- phide. Plumbites and Plumbates. Lead hydroxide, Pb(OH) 2 , re- acts with alkalies, giving plumbites, e. g., K 2 PbO 2 ; lead dioxide and alkalies give plunibates, e. g., K 2 PbO 3 . Normal plumbic acid would be H 4 PbO 4 (c/. silicic acid, 340). Its lead salt is Pb 2 PbO 4 , i. e., red lead. B. Tin (Atomic Mass, 119.) 454. Occurrence and Preparation of Tin. The only mineral abundant enough to serve as a source of tin is cassiterite, or tin-stone, SnO 9 . This occurs in Corn- COMPOUNDS OF TIN. 403 wall (England), in Australia, in the island of Banca, and in the Black Hills. Tin was known in very early times. Cassiterides was an ancient name for the Scilly Islands, owing to the fact that tin- stone was found there. Although tin has been carried away from Cornwall since the times of the Phoenicians, the mines there are still producing it. The metallurgy of tin consists, first, in roasting the ore, so as to oxidize and remove arsenic and sulphur. The tin oxide is then reduced with coal in a furnace, the metal being drawn off and cast into bars. These bars of impure tin are then slowly heated on a sloping hearth. The tin melts and runs down the hearth, leaving the unmelted impurities behind. 455. Properties and Uses. Tin is a white metal having a brilliant luster. It does not lose its luster in the air. It is soft and malleable, and melts at about 227 C. Tin reacts with hydrochloric acid, giving stannous chloride, SnCl 2 . With concentrated sulphuric acid it gives stannous sulphate, SnSO 4 , and sulphurous acid. Nitric acid oxidizes it to metastannic acid (H 2 SnO 3 ). The chief use of tin is to coat sheet-iron ; in this way the tin-plate of commerce is formed. It is also used to protect other metals, e. (/., copper and lead ; and in making alloys, e. (/., soft solder, pewter, bronze, bell-metal, etc. 456. Compounds of Tin. Tin forms stannous and gtannw compounds. Examples of the former are : stan- 404 LEAD, TIN, AND PLATINUM. nous chloride (SnCl 2 ), stannous oxide (SnO), and stan- nous sulphide, SnS. The corresponding stannic coin- pounds have the formulas SnCl 4 , SnO 2 , and SnS 2 . Stannic acid, like metastannic acid, is H 2 SnO 3 . Stannous chloride is easily oxidized to stannic chloride, and is, therefore, a good reducing agent. Thus, it reduces mercuric chloride to mercurous chloride, and even to mercury. (1) 2 HgCl 2 + SnCl 2 = SnCl 4 -f 2 HgCl. (2) 2 HgCl + SnCl 2 = SnCl, -f 2 Hg. Stannic chloride is a liquid. It is made by heating tin in chlorine. A solution of it is obtained by treating tin with aqua regia. Stannic oxide is made when stannic acid is heated, and when tin is burned in the air. This oxide shows its acid character, and its analogy to carbon dioxide, silicon dioxide, and lead di- oxide, by reacting with molten alkalies to form stannates, e. g., sodium stannate, Ka 2 SnO 3 . Stannous sulphide is a brown powder, formed when tin-foil is heated with sulphur, and when hydrogen sulphide is passed into the solution of a stannous salt. Stannic sulphide separates as a yellow precipitate when hy- drogen sulphide is passed into a stannic salt solution. Both sulphides react with alkaline sulphides, forming sul- phostannates, which are soluble in water (cf 314 and 318). C. Platinum (Atomic Mass, 195). 457. Occurrence and Preparation. Platinum is found native in a few places, chiefly in western Siberia. Native platinum is usually mixed with five other rare metals, all belonging to the eighth periodic group. These are : CHLORPLATINIC ACID. 405 palladium, ruthenium, rhodium, osmium, and iridium. About 75% of the ore is platinum. The ore is treated with aqua regia, which reacts with the platinum and some iridium. The resulting chlor- platinic acid, H 2 PtCl 6 (cf. bottom of page), is treated with ammonium chloride, producing a precipitate of ammonium Morplatinate, (NH 4 ) 2 PtCl 6 . When this is heated strongly, metallic platinum results. The small quantity of iridium is not removed. 458. Properties and Uses. Platinum is a grayish- white metal, over 21 times as heavy as water. An- nas no action upon it ; and the temperature of the oxy- hydrogen flame is needed to melt it. Platinum is not attacked by the common acids ; but it reacts with aqua regia and with chlorine- and bromine- water. Fused alkalies also act upon it. The resistance of platinum to most chemicals and the high temperature at which it fuses make it very useful in the laboratory. It is used in the form of foil, cru- cibles, wire, and other utensils. Large retorts of plati- num are used in concentrating and distilling sulphuric acid (cf. 185). 459. Chlorplatinic Acid. When platinum is treated with aqua regia, the platinum chloride (PtCl 4 ) formed unites with hydrochloric acid, giving chlorplatinic acid, H 2 PtCl e . The solution of this substance gives with potassium salts and with ammonium salts precipitates of potassium chlorplatinate, K 2 PtCl 6 , and of ammonium chlorplatinate, (NH 4 ) 2 PtCl 6 , re- spectively. The corresponding sodium salt is soluble. LABORATORY DIRECTIONS. (For the Student.') 1. Provide yourself with an apron and a pair of sleeves (rubber is the best material for these) ; also with soap and towel, and a white cloth about a yard square. The cloth is to be used for wiping apparatus. 2. Work by yourself ; and give your own descriptions, ob- servations, and calculations, not those of another. 3. Record at once all the observations you make in connec- tion with an experiment. See that your notes contain the answer to every question, direct or implied, that occurs in the laboratory exercise. Write neatly and distinctly. If the notes of two experiments occur on the same page, separate them by at least two centimeters of space. 4. Have a place for everything. Throw away nothing until you are sure you are through with it. Throw nothing but liquids into the sink. Put other waste materials into the proper receptacle. 5. If an experiment is unsatisfactory, repeat it until you are successful ; but first learn the probable cause of your error. 6. When you enter the laboratory, examine your table, and see that everything has been left as it should be by the persons who share the table with you. If anything is wrong, report the fact at once to the instructor. When you leave, see that the water and the gas are turned off, and that everything on your table is in good order. LABORATORY EXERCISES. EXPERIMENT I. THE BUNSEN BURNER. Apparatus. Bunsen burner, test tube, test-tube holder (see note below). Materials. Matches, water. a. Examine carefully the Bunsen burner on your desk. Take it apart, and draw a sketch of each part. b. Put the burner together, close the holes at the base, and connect with gas supply. To light the burner, turn on the gas and then hold a lighted match near the side of the burner and about one- half a centimeter below its mouth. Note the character of the flame ; is it luminous or not ? Now open the holes care- fully until therluminous region has just disappeared. This is the " Bunsen " flame. For most work it should be 7 to 10 centimeters (3 to 4 inches) high. The holes of the burner should be open far enough to prevent a deposit of soot upon the object heated, but not far enough to cause the flame to makfe. a noise. c. Introduce quickly into the center of the Bunsen flame, one-half a centimeter above the burner, the head end of a match. Result ? Is the gas in this region burning ? l 3 LABORATORY EXERCISES. To heat an object effectively, place it higher up in the flame / the best place is just above the apex of the dark, inner cone of unburned gas. Locate this region. d. Put 5 c.c. water into a test tube, and make a note of the height of the column of water in centimeters. When- ever you are asked to take 2, 5, 10, etc., cubic centimeters of anything, refer to this experiment, and use the length of the column just measured as your unit. e. Heat the water in the test tube to boiling. To do this properly have the outside of the tube dry ; hold the tube in the holder, and incline the tube at an angle of about 45 to the table top. Then introduce the bottom of the tube into the effective region (cf. c) of the flame. Heat only the part of the tube containing the liquid / if the flame strikes the glass above the liquid level, the tube may crack. Do not hold the tube still, but move it gently in the flame. When boiling begins, raise the tube a little above the flame, always keeping it inclined, so that the water may not " boil over." /. These directions are general, and will apply whenever you heat liquids in test tubes. Note. A very convenient test-tube holder can be made by folding a piece of writing paper twice, so as to produce a strip about 1 cm. wide and 10 to 15 cm. long. This is placed about the tube like a holder. The free ends are held together close to the tube. CUTTING AND BENDING GLASS TUBING. 3 EXPERIMENT II. CUTTING AND BENDING GLASS TUBING. Apparatus. Bunsen burner, "wing-top" or illuminating gas burner, file. Materials. Piece of soft glass tubing more than 15 cm. long. a. Cut off a piece of glass tubing 15 cm. long. To do this, make on the tubing a file mark in a plane perpendicular to the length of the tubing ; grasp the tube in both hands, and place the thumb nails together opposite the scratch. By pushing gently with the thumbs and at the same time pulling with the hands you will succeed in breaking the tubing so that the ends are fairly regular. b. Round off both ends of the 15 cm. tube by turning them about in the proper region of the Bunsen flame until the edges become red hot. Let the ends cool. c. Bend the 15 cm. tube at its middle into the form of a right angle. For this purpose use a flat Bunsen flame produced by a " wing- top " attachment or a flat illumi- nating flame. . Take the tube in both hands, one at each end, and hold its central part lengthwise with and over the flat flame. At the same time twirl the tube between thumbs and fore- fingers. Then lower the tube keep turning it into the upper part of the flame, and heat until you find that the glass is fairly soft. Then bend gently to a right angle. d. If you used the Bunsen flame, anneal the glass at the bend by closing the holes of the burner and allowing the hot LABORATORY glass to cool first in the smoky flame. When the bend is covered with soot, support it so that it will not touch a cold object. When the tube is cold, wipe off the soot. EXPERIMENT III. EFFECT OF HEAT UPON "RED PRECIPITATE." Apparatus. Small ignition tube of hard glass, rubber con- necting tube, delivery tube, pneumatic trough, test tube, ring stand, clamp. Materials. Pine splinter, red precipitate. a. In a small tube of hard glass sealed at one end and abou ; 10 cm. long " ignition tube " -place a layer of red precipitate not more than one-half a centimeter thick. In a basin containing water, invert a test tube of water. See that no ah* bubbles remain in the test tube. Vessels for holding water over which gases are collected are called " pneumatic troughs." Attach to the ignition tube by means of a piece of rubber tubing a delivery tube long enough to reach to the bottom of the pneumatic trough. Support the ignition and deliv- ery tubes so that the closed end of the ignition tube is only a little lower than its other end, and so that the red precipi- tate may be heated in the hot portion of the Bunsen flame. b. Begin to heat slowly, keeping the flame in motion. Note any change in color of the red precipitate. Afterward heat strongly with a steady flame until all of the powder disappears. Collect over water anything that escapes from the delivery tube by allowing it to displace the water of the SOLUTION, FILTRATION, AND EVAPORATION. 5 test tube. When the operation is over, remove the delivery tube from the water before removing the flame. Why ? c. Cover the mouth of the test tube under water with the thumb, remove tube from water, invert, and introduce a pine splinter with a spark on the end of it. Result ? Is the gas in the test tube air f d. When the ignition tube is cool, invert it and strike its open end sharply against the table. Result? What sub- stance is this ? On what part of the tube did it collect ? Why? e. If by a chemical change we mean one in which at least one new substance is formed, would you call this a chemical change, or not? EXPERIMENT IV. SOLUTION, FILTRATION, AND EVAPORATION. Apparatus. Glass rod 15 cm. long (unfinished), file, two beakers of about 50 c.c. capacity, ring stand, wire gauze, fun- nel, funnel support (small ring of ring stand), evaporating dish. Materials. Coarse salt, filter paper. a. Make a glass stirring rod 15 cm. long, cutting off a piece from a larger one, just as in Experiment II, a. Round off both ends in the flame. b. Put into a beaker about 20 c.c. cold water, add 5 grams salt, and heat the beaker over the flame until its contents boil. Before heating the beaker see that it is dry on the outside, then place it upon a wire gauze supported on the ring stand. Move the flame about under the ga-uze until the beaker has become warm / then put the burner under 6 LABORATORY EXERCISES. the center of the beaker. The height of the gauze above the burner should be so great that the bottom of the beaker may be a little above the apex of the dark inner region of the flame. Note. Always follow these directions when you are heating a beaker, an evaporating dish, or a flask, unless there is some special reason for not doing so. What becomes of the salt? Of the dirt? c. Next, filter the solution. You need a funnel, a support (see above), a filter, the glass rod made in a, and a second ' beaker. Fold the circular filter twice in lines at right angles to each other. Press the folded edges between thumb and forefinger, but not between the nails. Open the filter so that it shall form an inverted cone which just Jits the fun- nel. One-half of the conical surface is made up of three of the quarters into which the paper was folded ; the remain- ing quarter of the paper makes up the other half of the cone. d. Hold the filter in place in the funnel, and wet it com- pletely ; it should adhere everywhere to the inner surface of the funnel, and its point should extend a little into the stem of the funnel. Pour the salt solution down the glass rod to the filter. The glass rod should touch the lip of the beaker ; and the stem of the funnel should touch the side of the beaker beneath it. Always follow these directions in filtering an insoluble solid from a solution. e. Does anything remain on the filter ? We call it the residue. What passes through is the filtrate. HYDROGEN. 7 A substance which remains mixed with a liquid, but not dissolved in it, is said to be " suspended in," or " held in suspension by " the liquid. A suspended substance becomes, after filtration, a residue. f. Pour the filtrate of c into an evaporating dish, and heat (for precautions, cf. b) over the flame. Boil off the water until a solid begins to separate out ; then set the dish aside until it is cold, or until the next laboratory period. What is the solid obtained ? Ts this separation of the salt from the dirt a physical or a chemical operation? EXPERIMENT V. HYDROGEN. Apparatus. Generating flask, or bottle of 250 c.c. capacity, two-holed stopper, funnel tube, right-angled tube, rubber con- nector, delivery tube, pneumatic trough, squares of glass or of cardboard, two or more wide-mouth collecting bottles (250 c.c.). Materials. Zinc, dilute sulphuric acid (one part by volume of acid to four volumes of water), pine splinter, cupric sulphate solution. a. To a 250 c.c. flask containing enough zinc to cover the bottom fit a two-holed stopper. One of the holes is for a funnel tube reaching to within one-half a centimeter of the bottom of the flask when the stopper is in place ; the other hole contains a bent tube attached by a rubber connector to a delivery tube. The delivery tube reaches to a pneumatic trough containing two bottles filled with water and inverted. The level of the water in the trough should be about two 8 LABORATORY EXERCISES. centimeters (one inch) higher than the months of the in- verted bottles when the bottles are in place. b. To invert bottles in the trough without letting in air, fill them to overflowing with water,.cover their mouths with slips of glass or cardboard, press the latter against the bottle, and invert quickly under water. Then remove the covers. To remove a bottle full of gas from water, slip under the mouth of the bottle, under water, a glass or cardboard cover, and hold it in place as before. Leave a filled bottle with its mouth under water until used, if possible. Whether a bottle of gas shall be placed upright or in- verted upon the table depends upon the specific gravity of the gas. c. Caution. Keep all flames at least one meter (about three feet) away from the flask in which hydrogen is made. See that the stopper of the generating flask is tight, and add enough of the dilute sulphuric acid to immerse the lower end of the funnel tube. Tell what takes place in flask, funnel tube, and pneumatic trough. Explain each phenomenon. If action is not vigor- ous add a few drops of copper sulphate solution. Result ? If evolution of gas ceases or becomes slow before you are through, add more acid. The gas produced is hydrogen. d. Fill the two bottles with the gas and refill them after using. Reject the first bottleful collected by turning it mouth upward. Why not use it? Why turn it mouth upward ? Keep the second bottle inverted and introduce into its middle part a burning pine splinter 15 to 20 cm. long. Hold HYDROGEN. 9 the splinter steady 20 to 30 seconds. Result ? Does the gas burn ? Where ? Does the splinter continue to burn in the hydrogen? Is hydrogen combustible or a supporter of combustion ? Turn a third bottle of the gas mouth upward one minute, and repeat the test with the burning splinter. Results? From the result compare the specific gravity of hydrogen with that of air. e. Place the mouth of a fourth bottle of gas over the mouth of an upright bottle of air. Hold the bottles to- gether and reverse their positions. After one minute apply a lighted match to the lower bottle. Result ? To . the upper. Jesuit ? What conclusion as to the diffusibility of hydrogen ? f. Have a fifth (and last) bottle only half full of gas ; incline it, and then raise it slowly from the water so that air displaces the remaining water. Carry bottle, mouth down, to a flame. Result? Explain difference between this re- sult and the combustion of hydrogen free from air. g. From the experiment tell whether hydrogen is very soluble in water, or not. h. Pour the liquid and the unused zinc from the flask in- to a beaker. -' If the zinc has all dissolved, or if there seems to be enough acid to dissolve all of it, add more zinc. Leave until action ceases. i. Examine the beaker; has anything separated from solution? If so, re-dissolve it by heating the beaker on the wire gauze, and filter hot. ( Care /) Collect the filtrate in another beaker or an evaporating dish, and let it stand some hours. Result ? The substance you obtain is crystallized zinc sulphate. 10 LABORATORY EXERCISES. What two new products resulted from the action of zino and dilute sulphuric acid ? EXPERIMENT VI. EQUIVALENT OF MAGNESIUM. Apparatus. Balances, pneumatic trough, wide-mouth bot- tle (250 c.c.), graduated jar, glass or cardboard cover. Materials. Magnesium wire, dilute (5%) sulphuric acid. a: In a pneumatic trough containing water to the depth of about 3 cm. place a piece of magnesium wire the exact weight of which is known. There should be not more than 0.2 gram. b. Get the exact capacity in cubic centimeters of a wide- mouth bottle \)j filling it with water and pouring the water into a graduated vessel. The bottle should hold at least 250 c.c. c. Fill the bottle with 5% sulphuric acid, and invert it in the pneumatic trough as far from the magnesium as pos- sible. See that the bottle is free from air bubbles. Now slide the mouth of the bottle, under water, over the magnesium. Result ? d. When all the metal has disappeared, let the collected gas cool to room temperature for 5 minutes. Then add water of room temperature to the bowl, if necessary, so that the level of water in bottle and bowl shall be the same. Why ? Protect the bottle from the heat of the hand by grasping it with a towel ; then slip under its mouth a glass or card- OXYGEN. 11 board cover, and invert quickly ', so as to lose none of the water in the bottle. Bring a flame to the mouth of the bottle at once. Re- sult? The gas is hydrogen. The other product of the reaction is magnesium sulphate ; it remains in solution. e. Get the volume of the water remaining in the bottle by means of the graduated vessel. Then obtain by difference the volume of hydrogen. To get the weight of the hydrogen multiply its volume m cubic centimeters by the weight of 1 c.c. Get the weight of 1 c.c. under the conditions of the experiment from the teacher. What is your result? f. Solve the following proportion for x : weight of mag- nesium : weight of hydrogen :: x : 1. Result? x will be the equivalent of magnesium, i. e., the number of grams of magnesium required to liberate 1 gram of hydrogen (in this case from dilute sulphuric acid). EXPERIMENT VTI. OXYGEN. Apparatus. Mortar and pestle (?), test tubes, ring stand and clamp, one-holed stopper, delivery tube, pneumatic trough, 3 collecting bottles, glass or cardboard cover, deflagration spoon. Materials. Powdered potassium chlorate and manganese dioxide, pine splinter, sulphur, iron wire (picture cord) at least 15 cm. long. 12 LABORATORY EXERCISES. a. On a clean piece of writing paper mix carefully 6 grams powdered potassium chlorate with 5 grams powdered manganese dioxide. If the substances are not found in powdered form in the laboratory, grind them separately, in clean mortars, before mixing. b. Before you use the whole mixture, test the quality of a sample (1 c.c.) by heating it gently in an open test tube. If there is any evidence of violent combustion, or if large sparks appear, reject the mixture, and make a fresh one. A few small sparks indicate only traces of dust, etc. c. If the mixture is satisfactory, put it into a test tube supported by a clamp attached to a ring stand. The test tube is then fitted with a one- holed stopper and a delivery tube reaching under water in a pneumatic trough. Have 3 bottles filled with water and inverted in the trough. d. Heat the test tube gently from the top of the mixture downvmrd. Regulate the flame so as to keep the evolution of gas steady, but not violent. Keep the flame in motion, so as not to soften the glass. When the collecting bottles are full, first take the delivery tube out of the water, and then remove the flame. Why this precaution ? The gas is oxygen. e. Into one bottle of the gas put a glowing splinter as in Experiment III, b. Result ? Gradually lower the splinter into the bottle until combustion stops. What becomes of the splinter ? Of the oxygen ? To the contents of the bottle add 5 c.c. calcium hy- droxide solution (lime-water), cover with the hand, and shake vigorously. Result? KINDLING TEMPERATURE. 13 f. Note the odor of the gas in the second bottle. Then put into the bottle a deflagrating spoon containing burning sul- phur. Light the sulphur by holding the spoon in a flame. Have a cardboard cover with a small hole for the handle of the deflagrating spoon, and keep the bottle covered unti] combustion stops. Results? What becomes of the sulphur ? Of the oxygen ? Note the odor of the gas now in the bottle. Does this gas sup- port the combustion of a splinter ? Try it. g. Have the third bottle of oxygen covered and set up- right on the table. Draw aside the cover for a moment while you pour in 5 c.c. sand ; then replace the cover. Melt the sulphur left in the deflagrating spoon, and dip into it one end of a piece of iron picture cord. Light the sulphur tip, and at once hold the iron wire in the bottle of oxygen. Result ? Keep the wire in the gas until action ceases. Describe the product. Why was the iron tipped with sulphur? EXPERIMENT VIII. KINDLING TEMPERATURE. Apparatus. Wire gauze at least 15 cm. square, Bunsen burner, tongs. Material. Matches. a. Hold the wire gauze, by means of your tongs, 7 cm. above the Bunsen burner. Have the holes of the burner open as for the Bunsen flame. Now turn on the gas and bring a burning match from above down to the center of the gauze. Result ? 14 LABORATORY EXERCISES. Why does not the gas below the gauze take fire? Is there gas below the gauze ? Prove it. b. Let the gauze cool ; and then bring it down upon the Bunsen flame until the gauze is 6 to 7 cm. above the top of the burner. Result? Hold the gauze in place until it be- comes red hot. Result ? Explain. EXPERIMENT IX. ACTION OF SODIUM UPON WATER. Apparatus. Tongs, evaporating dish. Materials. Sodium, water, blue and red litmus paper, solid sodium hydroxide. Caution. Do not handle sodium with wet hands, or with wet forceps. Do not put sodium into the waste jar. On no account leave any sodium on or about your desk or in your locker. Sodium is usually kept under kerosene. a. What is the appearance of a freshly cut surface of sodium ? Is sodium hard or soft ? Heavy or light ? b. Hold a piece of sodium having a volume not greater than 8 to 10 c.mm. at ami's length by means of iron tongs or forceps, and drop it upon water in a small evaporating dish. Result? Apply a lighted match hold it at arm's length to the sodium while it is acting on the water. Result? c. After action has ceased, wet your fingers with the solu- tion, and rub them together. Result? If you get no de- cided result, add a second piece of sodium (dry hands) of the same size as the first, and repeat. WATER OF CRYSTALLIZATION. 15 d. Test the action of a drop of the solution upon a piece of blue litmus paper. Upon red litmus paper. Results ? e. Add a small piece (same size as sodium used) of so- dium hydroxide to 5 c.c. water. Results? Test solution with the fingers and with litmus. Results ? Compare re- sults with those in d. Conclusion. EXPERIMENT X. WATER OF CRYSTALLIZATION. Apparatus. Test tubes, iron saucer (sand bath). Materials. Crystals of zinc sulphate, of potash alum (po- tassium aluminum sulphate), and of cupric sulphate. a. Place a few crystals of zinc sulphate in a dry test tube, and warm gently. Results ? Is there evidence of water ? Where? b. Repeat , using a crystal of potash alum. Results ? c. Note the taste of another crystal of potash alum ; then heat it strongly in an iron dish until no further change occurs. Results ? When the ignited alum is cold, taste it. Result? Place it in 5 c.c. water in a test tube, and boil carefully for five minutes. When the water is cool, taste it. Result? Assuming that heat simply drove off crystal- water from the alum, upon what does the taste of crystalline alum seem to depend? d. Heat a crystal of copper sulphate (blue vitriol) strongly in an iron dish. Result ? When the residue is cold, add a few drops of water to it. Result ? Explain. 16 LABORATORY EXERCISES. EXPERIMENT XL EFFLORESCENCE. Apparatus. Evaporating dish. Materials. Crystallized sodium carbonate and sodium sul- phate (Glauber's salt). a. Expose a crystal or two of sodium carbonate to the air for at least twenty- four hours. Result ? b. Carefully weigh your evaporating dish, and then weigh into it accurately about 5 grams Glauber's salt (sodium sul- phate plus crystal- water). Let stand for at least twenty- four hours, and weigh again. Result? What change is there in the appearance of the substance? Record your results thus: Grams. (1) Weight of evaporating dish -j- Glauber's salt = (2) Weight of evaporating dish alone = (3) .-. Weight of Glauber's salt taken = (4) Weight of evaporating dish -f- residue = .*. gsfin or loss of water [subtract (4) from (1)] = EXPERIMENT XII. DELIQUESCENCE. Apparatus. Small beaker, evaporating dish or watch glass. Materials. Solid potassium hydroxide, granular calcium chloride. EFFECT OF TEMPERATURE. 17 a. In a small beaker place a piece of potassium hydroxide, and leave it exposed to the air at least an hour. Result ? b. Weigh an evaporating dish or a watch glass carefully, and then weigh into it accurately about 5 grams anhydrous calcium chloride. Let stand at least twenty-four hours, and weigh again. Results? Record the weighings as in Ex- periment XI, b. EXPERIMENT XIII. EFFECT OF TEMPERATURE ON SOLUTION. CRYSTALLIZATION. Apparatus. Beaker (50 c.c.), stirring rod. Materials. Potash alum, crystallized cupric sulphate (blue vitriol). a. Put 20 c.c. water into a beaker, add 10 grams pow- dered alum, and stir two minutes with the stirring rod. Does all the alum dissolve ? b. Heat the -beaker carefully on the wire gauze, stirring the contents. Result? Conclusion. c. Set the beaker with the hot solution in cold water, and stir rapidly until solution cools. Result ? d. Dry the outside of the beaker, and heat again as hi b. Result ? Let the solution stand undisturbed until it is cold. Result? Compare with c, and account for the difference. e. Repeat a, 5, c-, and 4 with 20 c.c. water and 15 grams powdered blue vitriol. Results ? 18 LABORATORY EXERCISES. EXPERIMENT XIV. PRECIPITATION. Apparatus. Test tubes. Materials. Solutions of lead nitrate, potassium chromate, barium chloride, and calcium sulphate. Dilute sulphuric acid ; alcohol. a. To 5 c.c. of lead nitrate solution in a test tube add an equal volume of potassium chromate solution. Result ? Let tube stand ten to fifteen minutes. Result? The pre- cipitate is lead chromate. b. Repeat a, putting together hot barium chloride solu- tion and dilute sulphuric acid. Result after ten to fifteen minutes ? The precipitate is barium sulphate. c. To 2 c.c. calcium sulphate solution add an equal volume of alcohol. Result? The precipitate is calcium sulphate. Note. The insoluble solids formed in a and b are not the only products of these reactions ; the other products are, how- ever, soluble. EXPERIMENT XV. CONSTANT PROPORTIONS. Apparatus. Evaporating dishes, beaker, balances, watch glass. Materials. Crystallized sodium carbonate, dilute hydro- chloric acid (1 volume concentrated acid to 1 volume water). a. Weigh your evaporating dish carefully, and then weigh into it accurately about 5 grams sodium carbonate crystals. Transfer the sodium carbonate without loss to a beaker cov- CONSTANT PROPORTION'S. 19 ere, A being removed. Then close the pinch- clamp. b. Into the test tube A put about 5 c.c of a mixture of equal parts of powdered, chemically pure potassium chlor- ate and manganese dioxide ; they must have been dried at B D FIG. 64. 56 LABORATORY EXERCISES. 120 C. for at least an hour. Get the weight of test tube and mixture accurately on the balances, and record it. c. See that the stopper is pressed, securely into the mouth of ./?, and then slip A carefully, but tightly, over its stopper. Now put about 50 c.c. water into C, raise C so that the water in and G are at the same level, open the pinch- clamp one minute, and then close it. Then put (J down on the table. Take the rubber tube carefully out of C and get the volume of the water in C / then pour the water back into C, and put the rubber tube in place. ISTow open the pinch- clamp, and hang it upon the glass tube at E. Do not allow the lower end of the rubber tube to get above the surface of the water in C. Why? d. Heat the mixture in A gently, beginning at the upper part of the mixture. The evolved gas forces water from B into C. When C is about half full, stop heating, and let A cool to room temperature. Then raise 15 or (7, as necessary, to make the water levels in both the same (be sure to keep the lower end of the rubber tube under water), close the rubber tube with the pinch- clamp, and get the volume of the water in C. This, minus the original volume, equals the volume of gas collected in 7?. Find the barometric height, subtract from it the cor- rection for the pressure of water vapor (see Appendix), and find, also, the temperature of the gas. Finally, weigh A and its contents accurately. e. Record your results thus : Grams. Weight of test tube -|- contents at first = Weight of test tube -j- contents afterward = .*. Weight of oxygen = Volume of oxygen at C. and mm. = c.c. .'. Volume of oxygen at C. and 760 mm. = c.c. Weight of oxygen obtained : weight of a liter at C. and 760 mm. :: volume (at C. and 760 mm.) : 1,000 c.c. BROMINE. 57 EXPERIMENT XLIIL BROMINE. Apparatus. Beaker, 100 c.c. flask, test tubes. Materials. Potassium bromide, powdered manganese diox- ide, dilute sulphuric acid, litmus paper, calico, carbon disul- phide, chlorine-water, sodium hydroxide. Caution. If possible, work in a gas-chamber or hood. a. Into a flask put an eighth of a test tube of potassium bromide (KBr) crystals, half as much powdered manganese dioxide, and half a test tube of dilute sulphuric acid. Sup- port the flask over wire gauze, and attach the cork stopper and a doubly bent delivery tube reaching into a test tube three-fourths full of cold water. The delivery tube must be without rubber connections. The test tube should rest in a beaker of water. b. Warm the flask carefully until a dark brown distillate passes over. Is it heavier or lighter than water ? Do not inhale the vapor, and do not get liquid bromine on your hands. When no more bromine comes over, remove first the de- livery tube and then the flame. The light-brown solution is " bromine- water." c. Wave the air from the test tube toward the nose. Odor of bromine ? Pour off as much bromine- water as pos- sible without pouring out the bromine, and add more water to the bromine. Pour a few drops of the bromine- water upon litmus paper and upon colored calico. Results ? d. To 3 c.c. water in a test tube add 1 c.c. carbon disul- 58 LABORATORY EXERCISES. phide, close tube with thumb, and shake vigorously. Re- sults? Where is the carbon disulphide? Now add c.c. bromine- water and shake again. Result to the color of the water ? To that of the carbon disulphide ? This effect on the carbon disulphide is a test for free bromine. e. To 5 c.c. of potassium bromide solution add 1 c.c. car- bon disulphide and shake. Result ? Now add two or three drops of chlorine-water (made as in Experiment XVI, f), close the tube, and shake it as before. Results ? Action of chlorine on potassium bromide ? Equation ? f. To the liquid bromine in the test tube add sodium hy- droxide solution, a c.c. at a time, shaking thoroughly. (Do not close the tube with the thumb !) Result ? The equation is, 2 NaOH + 2 Br NaBr -f NaBrO + H 2 O. sodium hypobromite EXPERIMENT XLIV. IODINE AND HYDRIODIC ACID. Apparatus. Test tubes, beaker, flask. Materials. Potassium iodide, manganese dioxide, sulphuric acid, iodine, carbon disulphide, chlorine- and bromine-water, starch, alcohol, hydrogen sulphide, silver nitrate solution, so- dium carbonate, litmus. a. Powder potassium iodide (KI), mix 1 c.c. of it with a c.c. of manganese dioxide, add 2 c.c. water and then 1 c.c. concentrated sulphuric acid. Result? When the IODINE AND HYDRIODIC ACID. 59 action slackens, warm the tube gently, and then let it cool. Describe what you find in the tube ? It is iodine. Com- pare its preparation with that of chlorine and bromine. b. Warm a crystal of iodine (gently, not to boiling) with 10 c.c. water for a few seconds. Does the iodine dissolve readily? Cool the water and add 3 c.c. of it to 1 c.c. car- bon disulphide. Shake the closed tube. Result ? This is a test for free iodine. Save the iodine solution. c. Shake 5 c.c. potassium iodide solution with 1 c.c. car- bon disulphide. Result? Add a drop of chlorine- water and shake again. Result ? What effect has chlorine upon potassium iodide ? Repeat, using bromine- water instead of chlorine- water. Write both equations. d. Make a starch solution as follows : Mix 2 c.c. powdered starch with 5 c.c. cold water, and pour the emulsion into 30 c.c. boiling water. Boil for a minute or two, and then cool. To 3 c.c. of the solution add a drop of the iodine solution of b, shaking. Result ? To 3 c.c. of the starch solution add one drop of a potas- sium iodide solution and then one drop of chlorine- or bromine- water. Result ? e. Heat a -crystal or two of iodine in a dry, inclined test tube. Result? Let cool. Result? Effect of iodine on the skin ? On wood and paper ? /. To the iodine of e add 5 c.c. ethyl alcohol, C 2 H 5 OH. In which is iodine more soluble, water or alcohol? An alcoholic solution is often called a tincture. g. To one-half a c.c. of powdered iodine in a flask add 20 c.c. water and then pass in hydrogen sulphide (gas-chamber !) until the iodine disappears. Results? Boil the solution gently two minutes, and filter it. Identify the precipitate 60 LABORATORY EXERCISES. by igniting a little on a piece of porcelain. Odor ? Test the filtrate with red and blue litmus. Results? Add a drop of it to 1 c.c. silver nitrate solution. Result? Add some to 1 c.c. solid sodium carbonate. Result ? What sub- stances are formed from hydrogen sulphide and iodine ? Equation ? EXPERIMENT XLV. COMPARISON OF THE HALOGEN ACIDS. Materials. Potassium chloride, bromide, and iodide; con- centrated sulphuric acid, litmus. a. Three test tubes have small amounts of potassium chloride, bromide, and iodide, respectively ; treat each with a few drops of concentrated sulphuric acid. Results ? Blow your breath across the mouth of each tube. Result ? Test the gas of each with blue litmus. Result ? Note carefully the odor of each gas. What odors beside that of the acid do you get in the tube of potassium iodide? Heat this tube. Result ? b. Which tube gives a colorless gas ? What colors the seas in each of the two other cases ? From the amount of O coloration, tell which of the three halogen acids is most easily decomposed into its elements. Which least. HYDROGEN PEROXIDE. 61 EXPERIMENT XLVI. HYDROGEN PEROXIDE. Materials. Hydrochloric acid, barium peroxide, starch so- lution, potassium iodide solution, manganese dioxide, potassium permanganate, ether, potassium dichromate solution, splinter. a. To 25 c.c. water add 5 c.c. concentrated hydrochloric acid and then 3 grams powdered barium peroxide, BaO 2 , a little at a time, stirring. Filter the solution ; it should con- tain Jiydrogen peroxide, H 2 O 2 . b. To 5 c.c. starch solution add a drop of potassium iodide solution, and then a few drops of the hydrogen peroxide solution. Result ? c. To 3 c.c, of the solution of a add 3 c.c. ether. Do they mix? Is the ether above or below? Now add one drop of potassium dichromate solution. Close tube and shake gently. Result ? d. To 5 c.c. of the hydrogen peroxide solution add 1 c.c. powdered manganese dioxide. Result ? Test gas with a glowing splinter. Result? e. To three crystals of potassium permanganate in a test tube add 2 c.c. water and then 5 c.c. of the hydrogen perox- ide solution. Result ? Test with glowing splinter. Result ? 62 LABORATORY EXERCISES. EXPERIMENT XLVII. PHOSPHORUS AND PHOSPHORIC ACID. Apparatus.- Test tubes, small ignition tube, tongs, evapo- rating dish, file. Materials. lied and yellow phosphorus, carbon disulphide, filter paper, phosphoric acid, ammonium hydroxide ; silver nitrate, disodium hydrogen phosphate, magnesium sulphate, ammonium chloride, and calcium chloride solutions. Caution. Ordinary, yellow phosphorus must be handled only with tongs, never with fingers ! It must be kept and cut under water. No pieces of it must get into your locker ; and every dish that has contained phosphorus must be heated, so that the phosphorus may be completely burned. Do not bring carbon diftulphide near a flame I a. Put half a c.c. of red phosphorus into a test tube, and add 3 c.c. carbon disulphide. Result? Filter, and let the carbon disulphide evaporate, without heating, in a hood, or where its vapor will not get near a flame. Result ? Was any phosphorus dissolved ? To 3 c.c. carbon disulphide add a piece of yellow phos- phorus not larger than a grain of wheat. Shake carefully a few minutes. Result ? Pour the solution, every drop of it, upon a piece of filter paper laid flat on a ring of the ring stand. Let the carbon disulphide evaporate without heating it. Result ? Rinse the test tube before putting it away. b. Into a small ignition tube put a layer of red phosphorus not more than 5 mm. thick, hold tube horizontal (tongs), and gently heat end containing the phosphorus. What collects ARSENIC. 63 on the cold part of the tube ? When the tube is cold, make a file-mark just below the deposit, and break the tube. Rub the deposit with a match stick. Result? Conclusion? Finally, heat both tubes red hot, so as to burn up all the phosphorus. Throw the pieces into iron or crockery jars. c. To 5 c.c. water add 1 c.c. concentrated (ortho) phos- phoric acid, neutralize in an evaporating dish (use litmus) with ammonia, and add silver nitrate solution. Result? The precipitate is silver orthophosphate, Ag 3 PO 4 . Write the two equations. Dissolve 2 c.c. powdered sodium hydrogen phosphate in 10 c.c. water. To half of the solution add calcium chloride solution. Result ? The product is secondary calcium phos- phate, CaHPO 4 . Equation ? d. To 5 c.c. magnesium sulphate solution add 1 c.c. am- monia-water and 1 c.c. ammonium chloride solution, and then the disodium hydrogen phosphate solution from c. Result ? The product is magnesium ammonium phosphate^ NH 4 MgP0 4 . 6H 2 0. (1) :Na 2 HPO 4 + NH 4 OH Na 2 ^H 4 PO 4 -f H 2 O. (2) ^a 2 NH 4 PO 4 -f- MgS0 4 MgNH^PO^ -f Na 2 SO 4 . EXPERIMENT XL VIII. ARSENIC. Apparatus. Small ignition tube, tongs, test tube, beaker. Materials. Arsenic trioxide (powdered), charcoal, hydro- chloric acid, hydrogen sulphide, ammonium sulphide, sodium hydroxide solution. 64 LABORATORY EXERCISES. a. Into a small ignition tube put powdered arsenic triox- ide, As 2 O 3 , to the depth of 5 mm. Hold the tube horizontal and at the side of the flame, so as^to heat only the end con- taining the powder. What happens? Now slip into the tube, almost to the arsenic trioxide, a piece of charcoal about 2 cm. long. Heat the charcoal red hot (have tube hori- zontal), and then incline the tube slightly so as to heat the arsenic trioxide while keeping the charcoal red hot. lie- suit ? Effect of charcoal upon the oxide ? Equation ? How does the oxide come into contact with the charcoal ? Sub- lime the arsenic obtained. b. Heat half a c.c. of arsenic trioxide with 8 c.c. dilute hydrochloric acid to gentle boiling. Result? Equation? Pour off from any undissolved material, and pass in hydro- gen sulphide for a minute. Result? If visible product is arsenic trisulphide, As 2 S 3 (its color?), write equation. Let settle, pour off supernatant liquid, and add 5 c.c. ammonium sulphide to residue, shaking. Result ? (CAUTION. Do not get ammonium sulphide on your hands ! ) The product now formed is ammonium sulpharsenite, (NH 4 ) 3 AsS 3 ; it is soluble. Treat solution with an excess of dilute hydro- chloric acid in a beaker. Result? c. Treat half a c.c. of arsenic trioxide with sodium hy droxide solution. Warm carefully. Result ? The solution contains sodium arsenite, Na 3 AsO 3 . From b and c would you say arsenic trioxide has acid, or basic, properties ? ANTIMONY. 65 EXPERIMENT XLIX. ANTIMONY. Apparatus. Mortar and pestle, funnel, ignition tube. Materials. Antimony, concentrated nitric and hydrochloric acids, hydrogen sulphide, ammonium sulphide, antimony triox- ide. a. What is the color of metallic antimony ? Is it heavy or light? Powder a small piece, and treat part of it in a test tube with concentrated nitric acid. Results ? b. Treat the remainder of the powdered antimony of a with 3 c.c. concentrated hydrochloric acid and 1 c.c. concen- trated nitric acid. Warm to start the action, if necessary. The solution contains antimony chloride, SbCl 3 . Let action continue for ten minutes ; then add 15 c.c. water. Filter, if necessary, and pass in hydrogen sulphide. If there is no action, dilute still more. Result? If the product has the formula Sb 2 S 3 , write the equation. Treat the antimony sulphide as you did arsenic trisulphide in Experiment XL VIII, b. c. Dissolve half a c.c. of tartar emetic in 5 c.c. water, add a drop of hydrochloric acid, and pass in hydrogen sulphide. Result ? Conclusion ? d. Heat antimony trioxide (Sb 2 O 3 ) in an ignition tube with charcoal, as you did arsenic trioxide. Results? 66 LABORATORY EXERCISES. EXPERIMENT L. BISMUTH. Apparatus. Mortar and pestle, beaker, test tubes. Materials. Bismuth, concentrated nitric and hydrochloric acids, bismuth nitrate crystals, hydrogen sulphide. a. What is the color of bismuth ? Is the metal heavy or light ? Malleable or brittle (test with the pestle) ? Treat a bit with concentrated nitric acid. Result ? Products ? b. To half a c.c. of bismuth nitrate crystals, Bi(NO 3 ) 3 , add 5 c.c. water, and shake. Result? If the product has the formula BiONO 3 , write the equation. Now add hy- drochloric acid (concentrated) a drop at a time, heating to boiling after each drop. Result? Use the least possible amount of acid. c. Put half of the solution from b into a beaker, and add much water. Result ? Compare with first part of b. d. To the remainder of the acidified solution of bismuth nitrate from b add hydrogen sulphide. Result ? The visible product is bismuth sulphide, Bi 2 S 3 . Write the equation. EXPERIMENT LI. BORAX AND BORIC ACID. Apparatus. Platinum wire sealed into glass rod, test tubes, beaker. Materials. -Borax, potassium dichromate, manganese diox- ide, hydrochloric acid, and sodium carbonate (solid). BORAX AND BORIC ACID. 6? a. Borax Bead. Make a loop 2 mm. in diameter on the end of a platinum wire sealed into a piece of glass tubing. Heat the loop to redness, and dip it into powdered borax, Na 2 B 4 O 7 . 10 H 2 O. Heat the adhering borax just within the outer edge of the Bunsen flame, at the place where the flame is widest. This is the fusing zone of the flame. What happens first? Heat until the borax melts to a transparent glass. If there is not enough borax to fill the loop, add more, and heat again. This glassy lump is called the borax bead. b. Touch the hot bead to a tiny speck (less than half as large as a pin's head) of potassium dichromate, K 2 O 2 O 7 , and heat at the top of the flame until the dichromate is com- pletely absorbed by the bead. Color? Remove the bead by plunging it while hot into water, and wipe it off the wire. c. Make a new bead, and touch it to a speck of man- ganese dioxide. Heat first in the oxidizing flame of the burner, i. e., just above the visible tip of the flame. Color? Now heat it in the reducing region, i. e., just above the tip of the bright blue interior cone. Heat it there persistently for five minutes, examining it from time to time. Result? Heat it again in the oxidizing flame. Result ? d. Boric Acid. Dissolve 5 grams powdered borax in 10 c.c. hot water, and add 10 c.c. concentrated hydrochloric acid. Set aside until next laboratory period. Result ? The product is boric acid, H 3 BO 3 . Filter off the crystals, wash them on the filter with a little cold water, and dry them on fresh filter paper. e. Dissolve the crystals of boric acid in hot water, and add the solution to a lump of sodium carbonate. Result ? C8 LABORATORY EXERCISES. EXPERIMENT LII. IONIZATION. Materials. Solutions of silver nitrate, potassium chloride, potassium chlorate, ammonia, and potassium ferrocyanide ; fer- rous sulphate crystals. a. Take 2 c.c. silver nitrate solution in each of two test tubes. To one tube add a few drops of potassium chloride solution. Name the precipitate from Experiment XVII, -j to /. Equation? b. To the second tube add potassium chlorate solution. Result? Is the same substance precipitated as in a? c. Powder a crystal of ferrous sulphate, FeSO 4 , and shake it with 5 c.c. water. Pour off the solution, and add to it a c.c. of ammonium hydroxide solution. Result ? If the pre- cipitate is ferrous hydroxide, Fe(OH) 2 , write the equation. d. To 3 c.c. of potassium ferrocyanide solution, K 4 Fe(CN) 6 , add ammonium hydroxide. Result? Is fer- rous hydroxide precipitated, as before ? e. Refer to Experiment XXXIX, a. Why does the ac- tion between sodium bicarbonate and tartaric acid take place only when water is present? EXPERIMENT LIIL HYDROLYSIS AND MASS ACTION. Materials. Antimony trichloride (crystalline or melted), hydrochloric acid. SODIUM COMPOUNDS. 69 a. To a small amount (half a c.c.) of antimony chloride, SbCl 3 , add 5 c.c. water, and shake. Result? The visible product is essentially antimony oxychloride, SbOCl, i. e., OH SbOH minus water. Write the equation. Compare the Cl result with Experiment L, #, where bismuth nitrate was used. b. To the precipitate add concentrated hydrochloric acid, a drop at a time, warming after each drop. Result? If the solution contains antimony chloride, SbCl 3 , write the equation. c. Add the solution obtained in b to 50 c.c. water. Re- sult? Add concentrated hydrochloric acid again. Result? d. Compare the equations of a and b. Write one of them, using, instead of the equality sign, the double arrow c *. In which direction does the reaction go chiefly when an excess of water is used? When an excess of acid is used? EXPERIMENT LIV. SODIUM COMPOUNDS. Apparatus. Test tubes, stopper and delivery tube, magni- fying glass, platinum wire, watch glass or glass slip. Materials. Sodium bicarbonate, lime-water, sodium car- bonate (solid and in solution), sodium chloride, calcium chlor- ide, barium chloride, sodium nitrate and sulphate, hydrochloric acid. a. Refer to Experiment IX for the properties of sodium and its action on water. 70 LABOEATOBY EXERCISES. b. Heat 2 c.c. powdered sodium bicarbonate carefully in a test tube having a delivery tube that passes into lime-water. Result? Is there any other volatile product? When no more gas is evolved (do not melt the test tube), let the product in the tube cool, and then add 2 c.c. cold water. Note the temperature effect. Compare with this the action of anhydrous sodium carbonate upon water. What are the products formed by heating sodium bicarbonate ? Equation ? c. Heat 2 c.c. sodium chloride with 5 c.c. water in a test tube; filter; and let some of the filtrate evaporate com- pletely on a glass slip or a watch glass. Examine the crystals with a magnifying glass, if possible. Their shape ? d. Dissolve a small piece of calcium, chloride, CaCl 2 , in 5 c.c. water, and add sodium carbonate solution. Result? Repeat, using barium chloride instead of calcium chloride. Result? Write the equations. e. Dip a platinum wire with a glass holder (cf. Experi- ment LI, a) into 5 c.c. concentrated hydrochloric acid in a test tube, and then heat the wire in the Bunsen flame until the flame remains colorless. If necessary, dip the wire more than once. Now wet the clean wire with the acid, dip it into powdered sodium chloride, and heat it. Effect on the flame? f. Clean the wire and repeat e, using sodium nitrate in- stead of sodium chloride. Repeat again with sodium sul- phate. What color do sodium salts give to the flame? POTASSIUM COMPOUNDS. 71 EXPERIMENT LV. POTASSIUM COMPOUNDS. Apparatus. Watch glass, iron dish, test tubes, beaker or evaporating dish, platinum wire, copper wire. Materials. Potassium chloride, sodium nitrate, sulphur, barium chloride solution, potassium hydrogen tartrate, lime- water, dilute sulphuric acid, concentrated hydrochloric acid, potassium nitrate, and potassium sulphate. a. Heat 8 grams of potassium, chloride and 10 grams of sodium nitrate with 20 grams of water until there is com- plete solution, and boil off half of the water over the wire gauze. Result? Let the precipitate settle and pour the solution into a test tube. Wash the residue with 5 c.c. cold water, and then dissolve it in the smallest possible amount of hot water. Pour a few drops of the solution in a watch glass and set aside. Result? Compare the crystals with those obtained in Experiment LIV, c. Conclusion ? What happens in the test tube containing the original solution ? The visible product is potassium nitrate, KNO 8 . . b. Mix 3 c.c. powdered potassium nitrate on a clean piece of paper with 1 c.c. powdered sulphur, and pour the mixture, at arm's length, upon a hot iron dish (use no wire gauze). Result? Let the product cool, boil it with 10 c.c. water in a test tube, and add to 5 c.c. of it barium chloride solution. Result ? See Experiment XXXYI, b. What is the product of the deflagration of potassium nitrate and sulphur ? c. Heat an iron dish red hot, and pour upon it 3 c.c. pow- dered potassium hydrogen tartrate, KHC 4 H 4 O 6 (cream of 72 LABORATORY EXERCISES. tartar). Results? Color of residue ? Heat it five minutes longer at red heat, pressing the mass down with a glass rod occasionally. When the dish is coql, treat the residue in a test tube with dilute sulphuric acid. After all evolution of gas ceases, identify the gas by placing in the mouth of the tube a looped copper wire holding a drop of lime-water. What remains undissolved? What substance would you find in plant ashes if the plants contained potassium salts of organic acids ? d. Clean a platinum wire as in Experiment LIY, e ; dip it into strong hydrochloric acid, and then into powdered potassium chloride, and heat it in the flame. Result ? Re- peat, using potassium nitrate instead of the chloride. Use the sulphate. Results? What color do potassium com- pounds give to the flame ? EXPERIMENT LVI. SOLUBILITY OF POTASSIUM CHLORIDE. Apparatus. Steam bath, water bath, or wire gauze ; evap- orating dish, balances. Materials. Powdered potassium chloride, distilled water. a. Make a saturated solution of potassium chloride by shaking 12 grams of the powdered substance in a clean flask with 25 c.c. distilled water at the temperature of the room. Continue shaking every little while for fifteen minutes. Record the temperature of the solution, and then weigh out accurately into your evaporating dish about 20 grams of the solution. Now evaporate (see b) the water until the residual AMMONIUM AMALGAM. 73 potassium chloride is perfectly dry, and get its weight. From the results calculate how much potassium chloride will dissolve in 100 grams of water at the room temperature. b. If possible, evaporate the solution of a on a steam or water bath. If this is impossible, evaporate slowly and care- fully on wire gauze, so as to avoid any loss by spattering. c. Record your results thus : Grams. Weight of evaporating dish -j- water -|- KC1 = Weight of evaporating dish alone = .'. Weight of water -f KC1 = Weight of evaporating dish -J- KC1 = Weight of evaporating dish alone = .'. Weight of KC1 = .'. Weight of water found : weight KC1 found :: 100 grams : x. EXPERIMENT LVII. AMMONIUM AMALGAM. DISTINCTIONS BETWEEN THE ALKALI METALS. Materials^ Ammonium chloride, sodium amalgam, sodium and potassium chlorides, tartaric acid, two unknown substances. a. Dissolve 2 c.c. ammonium chloride in 5 c.c. water, and add a piece of sodium amalgam (Na-J-Hg). Results? The product is ammonium amalgam. Note what happens to it. Odor ? Reaction of solution ? Note. Do not throw away the resulting mercury, but ask what to do with it. b. Add 5 c.c. water to 3 c.c. powdered potassium chloride 74 LABORATORY EXERCISES. and shake thoroughly. Pour off the solution and add to it 5 c.c. of a concentrated solution of tartaric acid, H 2 C 4 H 4 O 6 . Make this by shaking 5 c.c. powdered tartaric acid with 15 c.c. water. Wait for result. Result ? The product is po- tassium hydrogen tartrate. Equation ? c. Repeat #, using sodium chloride in place of potassium chloride. Result ? Repeat again, using ammonium chlor- ide in place of potassium chloride. Result ? d. From Experiment XXV, b and c, tell what happens when ammonium salts are treated with alkalies. How dis- tinguish between sodium salts on the one hand and ammon- ium and potassium salts on the other? Between sodium salts and potassium salts (two ways) ? e. Get from the instructor two unknown substances, and determine if they are salts of sodium, potassium, or ammo- nium. EXPERIMENT LVIII. CALCIUM. Apparatus. Triangle of iron wire, ring stand, blast-lamp, evaporating dish, platinum wire, and coin. Materials. Lumps of marble, lime-water, red litmus paper, old mortar, plaster of Paris, paper, calcium chloride, calcium sulphate, and ammonium carbonate solution. a. Touch a piece of wet red litmus paper with a piece of marble. Result? Support a lump of marble about 5 c.c. in volume on a triangle of iron wire laid upon a ring of the ring stand, and heat the marble in the hottest Bunsen flame in a blast-lamp, if possible. When the marble is cold, WATER OF CRYSTALLIZATION IN GYPSUM. 75 touch wet, red litmus with the part that was heated. Result ? Explain. What products are formed when marble is heated (cf. Experiment XL, b) ? Slake about 5 c.c. of quicklime by adding to it water, drop by drop, as long as the water is taken up readily. Wait for the result, and describe it. Is there a temperature effect ? Equation ? b. To a piece of old mortar in a test tube add dilute hy- drochloric acid. Identify the gas. What does fresh mortar absorb from the air ? c. Stir 10 c.c. plaster of Paris in an evaporating dish with enough water to form a fairly thick paste. Put the paste upon a piece of paper, and push into it a coin slightly coated with vaseline. At once wash the evapora- ting dish. Let the paste and coin stand an hour or more. Carefully remove the coin from the plaster. Result ? d. To a solution containing a calcium salt, i. e., calcium ionS) add ammonium carbonate solution. Result? See Experiment LIV, d. e. Clean a platinum wire as in Experiment LIV, e, and determine what color calcium chloride gives to the flame. Repeat with ^calcium sulphate. Be sure to have concen- trated hydrochloric acid upon the wire. EXPERIMENT LIX. WATER OF CRYSTALLIZATION IN GYPSUM. Apparatus. Evaporating dish, wire gauze, balances, Material. Powdered gypsum (not plaster of Paris). 76 LABOEATOEY EXEECISES. a. Weigh your evaporating dish (be sure it is clean and dry), and into it weigh accurately about 3 grams of finely powdered gypsum. Get the exact weight of the gypsum taken, and record it. b. Heat the evaporating dish on a clean wire gauze for ten minutes with the hottest Bunsen flame. Then let the dish cool, weigh it, and record the result. Now heat the dish again for five minutes, let it cool, and determine the weight. Compare the weight after the first ignition with that after the second. Keep your second weight as the final one. c. Record your results thus : Grams. Weight of evaporating dish -|- gypsum = Weight of evaporating dish alone = .". Weight of gypsum taken = Weight of evaporating dish -f- calcium sulphate = Weight of evaporating dish alone = .'. Weight of water found = .*. Per cent of water in gypsum = EXPERIMENT LX. STRONTIUM AND BARIUM. Apparatus. Platinum wire and test tubes. Materials. Strontium chloride and nitrate, barium chloride and nitrate ; solutions of strontium and barium chlorides ; ammonium carbonate solution ; dilute sulphuric acid. a. Treat 2 c.c. strontium chloride solution with a few drops of ammonium carbonate solution. Result ? Repeat, CRYSTAL-WATER IN BARIUM CHLORIDE. 77 using barium chloride in place of strontium chloride. Write equations. b. Treat 2 c.c. strontium chloride solution with dilute sulphuric acid. Result? See Experiment XXXVI, a. Equa- tion? c. Clean the platinum wire as in Experiment LIV, e, and heat a bit of strontium chloride in the flame. Repeat with strontium nitrate, Sr(NO 8 ) 2 . Results? d. Repeat c, using the corresponding barium salts. Re- sults ? How distinguish between calcium, strontium, and barium salts ? EXPERIMENT LXL WATER OF CRYSTALLIZATION IN BARIUM CHLORIDE. Apparatus. Evaporating dish, wire gauze, balances, air bath (?). Material. Barium chloride, chemically pure. a. Have your evaporating dish clean and dryland get its weight. Then weigh into it accurately about 3 grams of barium chloride ; this should be pure, dry, and finely powdered. b. Heat the evaporating dish with its contents in an air bath at 120 to 130 C. for one hour, then cool it ten minutes, and get its weight. Record your results as in Experiment LIX, c, and get the per cent of water in the crystallized barium chloride. 78 LABOEATOEY EXEECISEti. EXPERIMENT LXII. MAGNESIUM, Apparatus. Tongs, test tubes. Materials. Magnesium wire, dilute hydrochloric acid, solu- tions of magnesium sulphate and chloride, disodium hydrogen phosphate, and ammonium chloride and Iwdroxide, magnesite. a. Hold a piece of magnesium wire 2 cm. long in the flame (use tongs). Result'? Describe the product. b. Treat a second piece of magnesium with dilute hydro- chloric acid. Result? Identify the gas, and write the equation. See Experiment VI. c. To 2 c.c. of magnesium sulphate solution add sodium carbonate solution. Result? Repeat, using magnesium chloride instead of the sulphate. d. See Experiment XLVII, d, for the action of a solution containing a magnesium salt with disodium hydrogen phos- phate and ammonium hydroxide. Rewrite the equations here. Repeat that experiment with magnesium chloride solution instead of the sulphate. Equation ? e. Treat a small piece (half a c.c.) of magnesite with dilute nitric acid. Result ? Identify the gas, and write the equation. From Experiment XL, #, tell the effect of heat upon magnesite. ZINC. 79 EXPERIMENT LXIII. ZINC. Apparatus. File or sand-paper, knife, iron dish with flat bottom, test tubes, mouth blowpipe. Materials. Zinc, tin, lead, and copper ; zinc dust ; solutions of zinc sulphate, sodium hydroxide, and ammonium sulphide ; dilute sulphuric and hydrochloric acids ; hydrogen sulphide ; stick of charcoal. a. Clean part of a piece of zinc with a file or with sand- paper. Color? Is zinc hard or soft (use a knife or rough edge of glass) ? Place a burner below the center of an iron dish. At equal distances from the center place pieces of zinc, tin, lead, and copper, and determine the order in which they melt. Return the metals to the proper bottles. b. Heat a piece of zinc on charcoal with the oxidizing flame produced by the mouth blowpipe. Results ? To do this proceed as follows : Hollow out a depression near one end of the charcoal, and into it put the zinc. To make the blowpipe flame, have a luminous Bunsen flame 4 cm. high, and hold the blowpipe so that the flame produced will be inclined about 30 degrees to the horizontal plane. To make an oxidizing flame, hold the end of the blowpipe inside the luminous flame, a centimeter above the tip of the dark, inner cone. Hold the charcoal at such a distance that the zinc is in the outer, faintly-luminous part of the blowpipe flame. To make a reducing flame, hold the tip of the blowpipe just at the outer edge of the flame at its middle part, and hold the assay (here, zinc) much nearer the blowpipe than in the oxi- 80 LABORATORY EXERCISES. dizing flame. The proper region is just at the tip of the inner, light-blue cone of the blowpipe flame. c. What action has hydrochloric acid upon zinc? Equa- tion ? See Experiment V for the action of dilute sulphuric acid, and Experiment X for the behavior of zinc sulphate crystals, ZnSO 4 . 7 H 2 O, when heated. d. Mix 1 c.c. zinc dust with 5 c.c. sodium hydroxide so- lution, and heat carefully. Test evolved gas with a flame. Result? The solution contains sodium zincate, Na 2 ZnO 2 . Write the equation. e. To 2 c.c. zinc sulphate solution add a drop of sodium hydroxide solution. Result ? What, probably, is the pre- cipitate ? Equation ? Repeat with a second test tube. Now add to the first tube dilute hydrochloric acid, shaking. To the second tube add an excess of sodium hydroxide, shaking. Results? The alkaline solution contains sodium zincate. Equations ? What do these experiments show as to the nature of zinc hydroxide ? f. To 10 c.c. zinc sulphate solution add a drop of dilute sulphuric acid, and then hydrogen sulphide. Result? Tut the solution into a beaker and add 5 c.c. ammonium sul- phide solution, stirring. Result? The product is zinc sulphide, ZnS. Color? Equation? Add 10 c.c. water, stir the mixture, let it settle, and then pour off the supernatant liquid. Add 15 c.c. more water* stir, let settle, and decant , i. e., pour off the water. This is called " washing by decantation." To the zinc sulphide add dilute sulphuric acid. Result? What is the gas? Equation? Why was not zinc sulphide precipitated by hydrogen sulphide? CADMIUM. 81 EXPERIMENT LXIV. EQUIVALENT OF ZINC. Apparatus. Same as in Experiment VI. Materials. Zinc, in sheet form or in sticks ; dilute (5%) sul- phuric acid. a. Dissolve zinc in dilute sulphuric acid just as you did magnesium in Experiment VI, and find the volume of hy- drogen liberated by a known weight of zinc. Use from 0.45 gram to 0.55 gram of zinc. If the zinc is in sheet form, it will react readily; but a little impurity, chiefly carbon, will remain insoluble. If the zinc is pure, it will react with difficulty ; therefore wind about the zinc a piece of platinum wire or a narrow strip of platinum foil. Set the apparatus aside until the zinc is in solution ; then pro- ceed as in Experiment VI. b. Reduce the volume of hydrogen to standard condi- tions, and calculate the weight of the hydrogen obtained. Finally, solve for x in the proportion, Weight of zinc taken : weight of hydrogen obtained :: x: 1. The value of x will be the equivalent of zinc. EXPERIMENT LXV. CADMIUM. Maierials. ~ Cadmium sulphate, hydrogen sulphide, am- monium sulphide. 82 LABORATORY EXERCISES. a. Dissolve completely not more than 1 c.c. cadmium sulphate, CdSO 4 , in 5 c.c. water, and add hydrogen sulphide in excess. Result? The visible product is cadmium sul- phide^ CdS. Color? Equation? What other sulphides of the same color have you had ? Wash the precipitate by decantation, and treat it with 5 c.c. ammonium sulphide. Result? How distinguish between cadmium sulphide arid other sulphides of the same color ? EXPERIMENT LXYL MERCURY. Apparatus. Pipette (medicine dropper). Materials. Mercury, concentrated nitric acid, hydrogen sul- phide, hydrochloric acid, sodium hydroxide and potassium iodide solutions, ammonium hydroxide, zinc, and copper. Caution. Before working with mercury remove all rings. Do not throw mercury away ; but ask what you are to do with it. a. By means of a pipette take from the mercury bottle a globule three times as large as an ordinary water drop ; add to it 2 c.c. water and 2 c.c. concentrated nitric acid. Result ? Let stand until action stops ; this may take some hours. b. While waiting for a, dissolve a globule of mercury the size of a water drop in concentrated nitric acid ; this gives -mercuric nitrate, Hg(NO 3 ) 2 . Equation (cf. Experiment XXIX, g) ? Dilute with 15 c.c. water. c. To 2 c.c. mercuric nitrate solution (b) add hydrogen sulphide. Result? The precipitate is mercuric sulphide, HgS, Equation ? MEECUEY. 83 d. Add to separate portions of the nitrate solution, hydro- chloric acid, sodium hydroxide solution, and potassium iodide solution, respectively. Results? Add the potassium iodide drop by drop, noting changes. Write equations where possible. Note. With sodium hydroxide we should expect mercuric hydroxide, Hg(OH) 2 ; this, however, decomposes into the oxide and water. e. Note the result of a. The crystals are mercurous ni- trate, HgNO 3 ; pour out into a beaker, and add 15 c.c. water and a drop of strong nitric acid. /. To 2 c.c. of the mercurous nitrate solution of e add hy- drogen sulphide. The precipitate is mercuric sulphide and mercury. Write the equation. g. Repeat d with the mercurous instead of the mercuric nitrate. Results ? With sodium hydroxide the precipitate is mercurous oxide, Hg 2 O. Write the equations. Treat the precipitate produced by hydrochloric acid with ammonium hydroxide. Result ? h. Into the rest of the mercurous nitrate put a strip of zinc and a copper wire. Results? Now rub them dry. Results ? i. Classify the results of c, d, f, and g in five vertical columns. Formula of Precipitant. Mercuric Nitrate. Mercurous Nitrate. Formula of Ppt. Color. Formula of Ppt. Color. NaOH, etc. 84 LABORATORY EXERCISES. EXPERIMENT LXVIL COPPER. Apparatus. File or sand-paper. Materials. Copper wire, concentrated hydrochloric acid; solutions of cupric sulphate, ammonium hydroxide, sodium hy- droxide, and cupric nitrate ; grape-sugar ; iron nail. a. File a piece of copper bright. Color ? Is it hard or soft ? From Experiment LXIII give its fusibility compared with zinc, etc. By holding one end of the wire in the flame determine if it is a conductor of heat. b. From Experiments XXIX and XXXIV tell the action of nitric and sulphuric acids upon copper. Find out if cop- per reacts readily with concentrated hydrochloric acid. c. To 2 c.c. cupric sulphate solution add ammonium hy- droxide solution in excess. Result? Repeat with sodium hydroxide instead of ammonium hydroxide. Result? Re- peat, having the cupric sulphate hot, and then add the sodium hydroxide. Result? The last precipitate is cupric oxide, CuO. How formed (cf. Experiment LXVI, d and g) ? d. From Experiment XXXIII, b and c, tell the effect of hydrogen sulphide upon cupric sulphate. Equation? Pass hydrogen sulphide into cupric nitrate solution. Result? Equation? What is the effect of heating Hue vitriol (cf. Experiment X, d) ? e. Dissolve half a c.c. powdered grape-sugar, C 6 H 12 O 6 , in 2 c.c. water, and add it to 5 c.c. cupric sulphate solution. Now add sodium hydroxide solution, shaking until the precipitate first formed is redissolved, Color? Warm SILVER. 85 carefully, noting changes. Let stand. Results? Color of product? It is cuprous oxide, Cu 2 O. What effect had the grape- sugar ? f. Put an iron nail into cupric sulphate solution. Result ? EXPERIMENT LXVIIL SILVER. Materials. Silver foil, silver nitrate solution, nitric acid, sod- ium thiosulphate ; solutions of sodium chloride and potassium bromide, iodide, and cyanide ; filter paper; hydrogen sulphide. a. In a test tube treat a piece of silver foil with 2 c.c. concentrated nitric acid. Result ? Equation ? Dilute with water to 10 c.c. b. To 2 c.c. of the solution of a add 5 c.c. sodium chloride solution. Result ? Equation (cf. Experiment XVII, j) ? Boil the contents of the tube. Result ? Get the precipitate on filter paper, and expose it to sunlight. Result ? c. To 5 c.c. of solution a add 1 c.c. potassium bromide so- lution. Result? Heat to boiling, pour off the supernatant liquid, and add to half of the precipitate sodium thiosul- phate solution, Na 2 S 2 O3 (make this by dissolving the crystals in water). Result? Expose the other half on filter paper to sunlight. Result? d. To 1 c.c. silver nitrate solution add 1 c.c. potassium iodide solution. Result? Equation? e. To Ice. silver nitrate solution add hydrogen sulphide. Result ? Equation ? f. To 1 c.c. silver nitrate solution addpotassium cyanide 86 LABOEATOEY EXEECI8ES. solution, drop by drop. Result? Equation? Continue adding it, shaking, until it is in excess. Result? The so- lution contains the double cyanide, KCN.AgCN, i. e. y KAg(CN) 2 . Add sodium chloride solution. Result? Ex- plain the result in terms of the ionic theory (cf. Experi- ment LII). EXPERIMENT LXIX. ALUMINUM. Apparatus. Test tubes, tongs, blast-lamp. Materials. Aluminum wire and filings, white muslin, hy- drochloric acid ; solutions of sodium hydroxide, aluminum sul- phate, sodium carbonate, alum, ammonium hydroxide, and cochineal ; powdered alum, sodium bicarbonate, potassium sul- phate, ammonium sulphate, aluminum sulphate. a. Determine whether aluminum is a conductor of heat as in Experiment LXVII, a. Does the wire melt in the Bun- sen flame (use tongs) ? Try the blast-lamp. Result ? b. To 2 c.c. aluminum filings add 5 c.c. concentrated hy- drochloric acid, and warm. Result ? Test the gas. Equa- tion? c. Wash the filings remaining from #, by decantation, add 5 c.c. concentrated sodium hydroxide solution, and warm carefully. Determine the nature of the gas evolved. Re- sult? The solution contains sodium aluminate, Na 3 AlO 3 (cf. Experiment LXIII, d). Equation? d. To 5 c.c. of aluminum sulphate solution, A1 2 (SO 4 ) 3 , add 1 c.c. sodium hydroxide solution. Result? Equation? Get half of the precipitate into a second test tube, and add IKON. 87 an excess of sodium hydroxide solution. Result? If the solution now contains sodium aluminate, Na 3 AlO 3 , write the equation. To the other half of the precipitate add hy- drochloric acid. Result? Equation? Compare with this the behavior of zinc hydroxide. e. Dissolve as much ammonium sulphate as possible in 5 c.c hot water, and add to it in a beaker 5 c.c. water simi- larly saturated with aluminum sulphate. Cool the mixture. Result ? The product is ammonium alum. Heat again to complete solution, and let stand over night. Result ? Shape of crystals ? f. Repeat e, using potassium sulphate instead of am- monium sulphate. Results ? Compare the crystals. g. To 5 c.c. of the solution of any aluminum salt add sodium carbonate solution. Result ? Identify the escap- ing gas. The precipitate is aluminum hydroxide, A1(OH) 3 . Mix a cubic centimeter of powdered alum with a cubic centi- meter of sodium bicarbonate, and add water. Result ? h. To 1 c.c. of a solution of cochineal add 5 c.c. alum so- lution, immerse a piece of white muslin, and then add am- monium hydroxide solution, shaking. Results ? EXPERIMENT LXX. IRON. Apparatus. Test tubes, tongs, blast-lamp, magnet, beaker. Materials. Iron wire and filings, copper wire ; hydrochloric, sulphuric, and nitric acids ; solutions of potassium ferrocy- 88 LABORATORY EXERCISES. anide, ferricyanide, and sulphocyanate ; ammonia water, hy- drogen sulphide, ammonium sulphide, solid ferrous sulphate, and ferric chloride. a. Compare the heat conductivity of iron wire with that of copper. Test its magnetic properties ; its fusibility in the Bunsen flame and the blast- lamp. Results ? b. Treat 3 c.c. iron filings in a beaker with 20 c.c. dilute hydrochloric acid, stirring. Results ? Identify the gas. If the solution contains ferrous chloride, Fe01 2 , write the equa- tion. When action almost ceases, filter off 10 c.c. of the solution. Color of filtrate? c. Divide the filtrate of b into four parts. To the first add a few drops of potassium ferricyanide solution, K,Fe(CN) 6 . Result? This is "Turn&ulFs Hue? To the second portion add ammonia- water. Result? Equation? Note any change on standing in the air. To the third part add potassium f err ocy anide, K 4 Fe(CN) 6 . Result? To the last portion add potassium sulphocyanate solution, KSCN", Result ? Wash out your test tubes and beakers at once. d. Filter the remainder of the ferrous chloride solution ol #, and add 2 c.c. concentrated nitric acid. Heat carefully for two minutes in a beaker. Resulting color ? The solution contains ferric chloride and nitrate. To a drop of it in a test tube add a drop of potassium ferricyanide solution ; if it still gives a blue precipitate, add 2 c.c. more nitric acid, and boil again. Treat the resulting substance in four test tubes with the reagents used in c. Result in each case ? The precipitate from potassium ferrocyanide and a ferric salt is "Prussian blue." IE ON. e. Classify the results of c and d (last part) in five ver- tical columns. Formula of Keagent. Ferrous Chloride. Ferric Chloride. Precipitate or Solution? Color. Precipitate or Solution? Color. f. In a test tube shake 2 c.c. powdered ferrous sulphate with 10 c.c. water, pour off half of the solution, and pass hydrogen sulphide into it. Result? Does all the iron appear to be precipitated? Write the equation represent- ing the reaction you would expect to take place. From Experiment XXXIII tell the effect of dilute sul- phuric acid upon ferrous sulphide. Write the equation here. Compare these two equations. Conclusion? To the other half of the ferrous sulphate solution add five drops of dilute sulphuric acid, and pass in hydrogen sulphide. Compare with the result without the acid ? Now add am- monium sulphide. Result ? Equation ? g. Dissolve 1 c.c. ferric chloride, FeCl 3 , in 10 c.c. water, and pass in hydrogen sulphide at least two minutes. Re- sult? Boil the contents of the tube, and then filter. Test the filtrate with a drop of potassium ferricyanide solution. Result and conclusion ? Determine the nature of the residue on the filter paper by collecting it on a piece of porcelain and igniting it. Odor ? Write the equation for the action of hydrogen sulphide on ferric chloride. 90 LABORATORY EXERCISES. EXPERIMENT LXXL NICKEL AND COBALT. Apparatus. Platinum wire, test tubes. Materials. Nickel and cobalt and their nitrates ; solutions of the nitrates ; borax, sodium hydroxide solution, concen- trated, chemically pure hydrochloric acid. a. Give the physical properties of cobalt and nickel from an examination of the metals. Effect of a magnet ? b. To 2 c.c. nickel nitrate solution, Ni(NO 3 ) 2 , add a drop of hydrochloric acid and then hydrogen sulphide. Result ? Now add ammonium sulphide. Result ? Equation ? Ex- plain the results from Experiments XXXIII and LXX, f. c. Make a borax bead as in Experiment LI, a and ft, and determine the color given to it by nickel nitrate. d. Repeat b and c with cobalt nitrate, Co(NO 3 ) 2 , instead of nickel nitrate. Results ? e. To 2 c.c. cobalt nitrate solution add sodium hydroxide solution, a drop at a time, until it is in excess. Results ? f. To 2 c.c. cobalt nitrate solution add 5 c.c. concentrated, chemically pure hydrochloric acid. Result? Dilute with water. Result ? EXPERIMENT LXXII. MANGANESE COMPOUNDS. Apparatus. Platinum wire, test tubes. Materials. Manganese sulphate, potassium permanganate, CHROMIUM COMPOUNDS. 91 ferrous sulphate, grape-sugar, ammonia-water, hydrogen sul- phide, and ammonium sulphide. a. Dissolve 1 c.c. powdered manganese sulphate, MnSO 4 , in 5 c.c. water. To half of it add a drop of dilute sul- phuric acid and then hydrogen sulphide. Result? Now add ammonium sulphide. Result ? Color ? Equation ? Explain the results. b. To the other half of solution a add ammonia-water. Result ? Eq nation ? c. To 2 c.c. ferrous sulphate solution (cf. Experiment LXX, /) add potassium permanganate solution. Result? Continue, drop by drop, until the solution is just faintly pink. Now add ammonia-water. State and explain the result (cf. Experiment XXXIV, /). d. Dissolve a crystal of potassium permanganate, KMnO 4 , in w r ater, and add grape-sugar solution. Result? Explain (cf. Experiment LXVII, e). e. From Experiment XVI tell the action of manganese dioxide with hydrochloric acid ; from Experiment VII, with potassium chlorate ; from Experiment XL VI, with hydro- gen peroxide / and from Experiment LI, a, tell the color of the manganese bead. EXPERIMENT LXXIII. CHROMIUM COMPOUNDS. Apparatus. Chlorine generator, platinum wire, evaporating dish, test tubes. Materials. Solutions of potassium chromate and dichrom- 92 LABOR A TOE Y EXERCISES. ate, of chromic chloride, of potassium and sodium hydroxides; hydrochloric acid, alcohol, borax, barium chloride solution, chrome-alum. a. What is the color of solutions of potassium di- chromate (K 2 Cr 2 O 7 ), potassium chromate (K 2 CrO 4 ), and of chromic chloride (Cr01 3 ) ? b. Treat 1 c.c. of potassium dichromate solution with a drop of potassium hydroxide solution. Result? From the color tell what is formed. Complete the equation, K 2 Cr 2 7 + 2 KOH ? + ? c. To 1 c.c. potassium chromate solution add a drop of concentrated hydrochloric acid. Result ? What is formed ? 2 K 2 CrO 4 + 2 HC1 ? + ? -f ? How can a dichromate be changed to a chromate ? A chromate to a dichromate ? d. To 2 c.c. potassium chromate solution add barium chloride solution. Result? Equation? Repeat with chromic chloride instead of the chromate. Result? e. To 1 c.c. chromic chloride solution add a drop of sodium hydroxide solution. Result? Equation? Now add the alkali in excess, shaking. Result? The solution contains a chromite, NaCrO 2 . What other elements behave in this way ? See Experiments LXIII and LXIX. Save for g. f. Repeat e with potassium chromate instead of chromic chloride. Result ? In what three ways can a chromic salt be distinguished from a chromate? g. To the clear solution of e add chlorine gas until there LEAD. 93 is no further change. Do this in a gas-chamber. Results ? Test the resulting solution as in b, c, and d. Results ? How can a chromic salt be changed into a chromate ? A. To 10 c.c. potassium dichromate solution in an evapo- rating dish add 2 c.c. concentrated hydrochloric acid and 2 c.c. ethyl alcohol. Boil until bright green, but not to dry- ness. Test a part of the liquid with barium chloride solu- tion. Result ? With potassium hydroxide solution. What does the green solution contain? How can a chromate be changed to a chromic salt? See, also, Experiment XXXIV, / i. Refer to Experiment LI for the borax bead test. Re- peat with a tiny piece of chrome-alum. Result ? EXPERIMENT LXXXIV. LEAD. Apparatus. File or knife, test tubes, mouth blowpipe. Materials. Lead ; hydrochloric, nitric, and sulphuric acids ; lead nitrate, solutions of potassium chromate and sodium hy- droxide, lead oxide, stick of charcoal. a. File or cut off the coating on lead. Is it hard or soft ? Color? Try to mark on paper with lead. Result? Refer to Experiment LXIII, a, for its fusibility. Treat a small piece with hydrochloric acid, both the dilute and the strong. Results ? Wash the lead, and add 2 c.c. concentrated nitric acid and 2 c.c. water. Heat gently. Result? Write the equation (cf. Experiment XXIX). b. Heat one-fourth of a c.c. of lead monoxide on charcoal 94 L AS OH AT OE Y EXERCISES. in the reducing flame (mouth blowpipe). See Experiment LXIII, b. Result? How identify the product ? c. Dissolve 2 c.c. powdered lead nitrate, Pb(NO 3 ) 2 , in 15 c.c. water, heating. Cool, and add to 2 c.c. of the solu- tion 5 c.c. dilute hydrochloric acid. Result? Equation? Wash the precipitate by decantation, and heat it with 10 c.c. water. Result ? Cool the solution. Result ? d. To 2 c.c. of the lead nitrate solution add dilute sul- phuric acid. Result? Use potassium chr ornate solution instead of sulphuric acid. Result ? Equation in each case ? From Experiment XXXIII, d, tell effect of hydrogen sulphide upon lead nitrate. For the reduction of lead oxide by charcoal, see Experiment XL, a. e. To 2 c.c. lead nitrate solution add a drop of sodium hydroxide solution. Result? Equation? Now add an excess, shaking. Result? What three other hydroxides behaved in the same way ? See Experiment LXXIII, e. f. Put into the remainder of the lead nitrate solution a strip of zinc. Leave it at least ten minutes. Result? Equation (cf. Experiment LXVII, /) ? EXPERIMENT LXXV. TIN. Apparatus. Test tubes, stopper and deliver}^ tube, mouth blowpipe. Materials. Tin (granular and in a bar) ; concentrated hy- drochloric and nitric acids ; solutions of mercuric chloride, stannic chloride, and sodium hydroxide ; ammonium sulphide, hydrogen sulphide, zinc, sulphur, stick of charcoal. TIN. 95 a. Treat about 2 c.c. of small bits of tin with 10 c.c. con- centrated hydrochloric acid in a test tube. Warm gently to start the action, and when the effervescence is vigorous attach a stopper and delivery tube and collect the gas over water. Identify the gas. Result? The solution contains stannous chloride, SnCl 2 . Equation ? Let the action con- tinue at least ten minutes. b. From Experiment LXIII, a, compare the fusibility of tin with that of lead, etc. Hold a bar of tin near your ear, and bend it. Result ? What color has bright tin ? Is it hard or soft ? c. To 1 c.c. mercuric chloride solution, HgCl 2 , add 4 or 5 c.c. of your stannous chloride solution, and then heat. Note all the changes. The solution contains stannic chloride, SnCl 4 . Equation ? d. To 2 c.c. stannous chloride solution add 5 c.c. water and then hydrogen sulphide. Result? Color? Equation? Wash the precipitate by decantation, and add 5 c.c. am- monium sulphide (use an evaporating dish or beaker) and a small lump of sulphur. Warm gently, and stir. Result? Cool, and add dilute hydrochloric acid in excess. Result ? Compare the color with that of the original precipitate. e. To 2 cX5. stannous chloride solution add 1 c.c. concen- trated nitric acid, and heat gently. The solution contains stannic chloride. Dilute with 5 c.c. water, and pass in hydro- gen sulphide. Result? Color? Stannic sulphide is SnS 2 . Equation? Wash the precipitate by decantation, add am- monium sulphide and a bit of sulphur, and warm gently. Result? Add an excess of dilute hydrochloric acid. Re- sult? Compare with the color of the original precipitate, and with that obtained at the end of d. Conclusion ? 96 LABORATORY EXERCISES. f. To 2 c.c. stannic chloride solution add sodium hydrox- ide solution, drop by drop. Result? Add an excess. Re- sult? What other hydroxides have behaved in the same way? See Experiment LXXIV, e. g. Pour the solution of a from any unused tin, and put into it a strip of zinc. Result ? Equation ? Compare with Experiment LXXIV,/ h. Heat a piece of tin on charcoal in the oxidizing flame (mouth blowpipe). See Experiment LXIII, b. Results? APPENDIX. TABLE OF EQUIVALENTS IN ENGLISH AND METRIC UNITS. A. LENGTH. 1 centimeter 0.3937 in. 1 decimeter 10 cm. 1 meter = 100 cm. 1,000 mm. 1 meter = 39.37 in. = 3.28 ft. 1 kilometer =1,000 m. = 0.6214 mile. 1 inch =2.54 cm. 1 foot = 0.3048m. 1 mile =1.6094 km. B. AREA. 1 sq. cm. = 0.155 sq. in. 1 sq. m. = 10.764 sq. ft. =1.196 sq. yd. 1 sq. km. =0.385 sq. mile. C. VOLUME. 1 cu. cm. = 0.061 cu. in. 1 cu. m/' = 35.315 cu. ft. 1 liter = 1,000 cu. cm. = 1. 0567 qt. (U. S.) D. WEIGHT. 1 gram = 15.4324 grains. 1 kilogram =1,000 grams =2. 2046 Ibs. 1 grain = 0.0648 gram. 1 ounce (avoirdupois) = 28. 35 grams. 1 ounce (troy) =31.1 grams. i 11 APPENDIX. TABLE OF ATQMIC MASSES AND SPECIFIC HEATS.. ATOMIC MASSES.* SPECIFIC Clarke. Richai'ds. ! 1 I , \ A O H = l. 1 O = 16. O=16. Aluminum . . . . . Al ... 26.9 27.1 27.1 0.214 Antimony . . . . . Sb ... 319.5 120.4 120.0 0.0508 Ar^on . . . A . . . . 39.0 39.96 39.92 Arsenic As . 74.45 75.0 75.0 0.0814 Barium .... . . . Ba ... 136.4 137.40 137.43 Bismuth . . . . . . Bi ... 206.5 208.1 208.0 0.0308 Boron .... . . B . . . . 10.9 11.0 11.0 0.3G6 Bromine . . . . . . Br ... 79.35 79.95 79.955 0.0843-f Cadmium . . . ... Cd ... 111.55 112.4 112.3 0.0567 Caesium . . . . . . Cs ... 131.9 132.9 132.9 Calcium .... . . . Ca ... 39.8 40.1 40.1 0.170 Carbon .... . . . C . . . . 11.9 12.0 12.001 0.459ft Cerium .... . . . Ce ... 138.0 139.0 140. 0.0448 Chlorine . . . . . . Cl ... 35.18 35.45 35.455 Chromium . . . . . Cr ... 51.7 52.1 52.14 0.100 Cobalt ... Co ... 58.55 59.00 59.00 0.107 Columbiuin . . . . Cb ... 93.0 93.7 94. Copper .... ... Cu ... 63.1 63.60 , 63.60 0.0952 Erbium .... . . . Er ... 164.7 166.0 166. Fluorine . . . . . . Fl ... 18.9 19.05 19.05 Gadolinium. . ... Gd ... 155.2 156.4 156. ? Gallium .... ... Ga ... 69.5 70.0 70.0 0.079 Germanium . . . . Ge ... 71.9 72.5 72.5 i Glucinum . . . . . Gl ... 9.0 9.1 9.1 0.058 Gold . . . Au . . . 195.7 197.2 197.3 0.0324 Helium .... . . . He . . . 3.93 3.96 3.96 Hydrogen . . . . . H. . . . 1.000 1.008 1.0075 . . . In ... 113.1 114.0 114. 0.0570 Iodine . . I . . 125.89 126.85 126.85 0.0541 Iridium .... . . . Ir. . . . 191.7 193.1 193.0 0.0326 Iron . Fe . 55.5 55.9 55.9 0.114 Krypton . . . . . . Kr ... 81.15 81.76 81.7 Lanthanum . . . . La ... 137.6 138.6 138.5 0.0449 Lead . . . Pb ... 205.36 206.92 206.92 0.0314 Lithium . . . . . . Li ... 6.97 7.03 7.03 0.941 Magnesium . . ... Mg ... 24.1 24.3 24.36 0.250 Manganese . . ... Mn ... 54.6 55.0 55.02 0.122 Mercury . . . ... Hg ... 198.50 200.0 200.0 0.03 I9f *Table of Atomic Masses, prepared by Prof. F. W. Clarke; " Journal of the American Chemical Society," Vol. XXIV, No. 3; March, 1902. t Solid, ft Diamond. APPENDIX. Ill TABLE OF ATOMIC MASSES AND SPECIFIC HEATS.- Continued. ATOMIC MASSES. SPECIFIC HEATS. 0.0722 0.108 0.0311 0.0593 0.189* 0.0324 0.166 0.0580 0.0611 0.0762f 0.203f 0.0570 0.293 o.nstt 0.0474 0.0335 0.0276 0.0562 0.1485 0.0334 0.0277 0.0955 0.0662 Molybdenum . . Neodymium . . Neon Nickel Nitrogen . . Mo . . . . . Nd . . . . . Ne . . . . . Ni ... N Clarke. H = l. O = 16. 95.3 96.0 142.5 143.6 19.8 19.94 58.25 58.70 13.93 14.04 189.6 191.0 15.88 16.000 106.2 107.0 30.75 31.0 193.4 194.9 38.82 39.11 139.4 140.5 102.2 103.0 84.75 85.4 100.9 101.7 149.2 ? 150.3 ? 43.8 44.1 78.6 79.2 28.2 28.4 107.11 107.92 22.88 23.05 86.95 87.60 31.83 32.07 181.5 182.8 126.5 127.7 158.8 160. 202.61 204.15 230.8 ? 232.6 ? 169.4 170.7 118.1 119.0 47.8 48.15 182.6 184. 237.8 239.6 51.0 51.4 127. 128.0 171.9 173.2 88.3 89.0 64.9 65.4 89.7 90.4 Richards. = 16. 96.0 143.6 19.94 58.70 14.04 190.8 16.00 106.5 31.0 195.2 39.14 140.5 103.0 85.44 101.7 150.0 44. 79.2 28.4 107.93 23.05 87.68 32.065 183. 127.5 ? 160. 204.15 233. 171. ? 119.0 48.17 184. 238.5 51.4 128. 173. 89.0 65.40 90.6 Osmium Oxvaren . . . Os ... . . . . . . Palladium . . . Phosphorus . . . . Pd ... . . P . . . . Platinum .... Potassium . . Pt ... K . Praseodymium . Rhodium .... Rubidium .... Ruthenium . . . Samarium . . . Scandium .... Selenium .... Silicon Silver Sodium Strontium . . . Sulphur . . Pr ... . . Rh . . . . . Rb . . . . . Ru ... .. Sin ... . . Sc ... . . Se ... .. Si .... . . Ag . . . . . Na . . . . . Sr ... . . S . . . . Tantalum .... Tellurium . . . Terbium .... Thallium .... Thorium .... Thulium .... Tin . . . . Ta ... . . Te . . . . . Tr ... . . Tl ... . . Th . . . . . Tin . . . Sn Titanium .... Tungsten .... Uranium .... Vanadium . . . Xenon Ytterbium . . . Yttrium .... . . Ti . . . . . . W. . . . . . Ur . . . . . V . . . . . . Xe . . . . . Yb . . . . . Y . . . . Zu Zirconium . . . . . Zr ... * Yellow. t Crystalline. ft Rhombic. IV APPENDIX. TENSION OF AQUEOUS VAPOR IN MM. OF MERCURY (REGNAULT). TEMP. TENSION. TEMP. TENSION. TEMP. TKNSION. 0C. 4.6 11 C. 9.8 21 C. 18.5 1 4.9 12 10.4 22 19.7 2 5.8 13 11.1 23 20.9 3 5.7 14 11.9 24 22.2 4 G.I 15 12.7 25 23. G 5 G.5 16 13.5 2G 25.0 6 7.0 17 14.4 27 2G.5 7 7.5 18 15.4 28 28.1 8 8.0 19 1G.3 29 29.8 9 8.5 20 17.4 30 31.6 10 9.1 ... TABLE OF SPECIFIC GRAVITIES (WATER = 1). Acetic ucicl * .... 1.053 Lead 11.35 Alcohol (ethvl)* . . . 0.794 Limestone 3.2 2.67 Lithium 59 6 72 1 74 5.7 Manganese . . . 7 2 to 8 9.8 Mercury \ 13.596 2.63 Nickel .... S 57 8.3 Nitric acid (cone )* 1.42 Carbon (gas) .... Carbon disulphide * . 1.8 1.27 Phosphorus .... Platinum .... 1.83 21 5 Chloroform * .... 1 5 Potassium . . . 865 Coal (anthracite) . . . 1.2G to 1.8 Silicon .... 2 49 Cobalt 8.8 Silver .... 10 57 8.9 Sodium . . . 97 Diamond 3.53 Sulphur 2 03 Ether* 0.72 Sulphuric acid . . 1 84 (ihos 2 6 to 3.6 Tin 7 29 Gold 193 Water at C. ... 0999 Hydrochloric acid (cone.)* 1.22 0918 4.07 C. . . , ,, 100 C. . 1.000 0958 Iodine 495 (sea) . . . 1 026 Iron ... ... 7 8 Zinc 6 9 to 7.5 * At 15 C. t At 0C. WEIGHT (IN GRAMS) OF A LITER OF THE DRY GAS AT C. AND 760 MM. Air 1.293 0.717 Ammonia . . . 0.762 Nitric oxide .... 1.34 Carbon dioxide ... 1 977 1.256 Carbon monoxide . . Chlorine 1.251 3 18 Nitrous oxide .... 1.97 1.429 Hydrochloric acid . . 1.61 0896 Sulphur dioxide . . . 2.87 0.806 Hydrogen sulphide . . 1.542 APPENDIX. TABLES. Table I. HEAT OF FORMATION AND HEAT or SOLUTION OF SOME SUBSTANCES IN KILOGRAM-CENTIGRADE UNITS. NAME. Formula. Heat of Formation. Heat of Solution in Water. Ozone Oo 30 Water (liquid) H 2 O 68 4 Hydrogen peroxide .... Hydrogen chloride H 2 2 HC1 45.2 22 17 3 Hydrogen bromide HBr 12 19 9 Hydrogen iodide HI 6 1 19 2 Hydrogen sulphide .... H 2 S SO 2 2.7 71 4.6 77 H 2 SO 4 193 1 178 NH 3 12 8.4 Nitrogen tetroxide N,O Q Nitrogen dioxide N0 2 4 7 7 Nitric oxide . NO 21 6 Carbon dioxide . C&-- 19 6 Carbon disulphide .... ^-CO,- 97 Carbon monoxide CO 29 Phosphorus, red, fromyellow form Potassium hydroxide .... Potassium carbonate .... Potassium nitrate KOH K 2 C0 3 KNO 3 27.3 103.2 281. 119 13.3 6.5 85 Sodium chloride NaCl 97 6 1.2 Sodium hydroxide NaOH 102 9 9 Sodium carbonate NaoCOo 272 6 5 6 Ammonium chloride .... Calcium hydroxide .... Magnesium hydroxide . . . Aluminum hydroxide .... Ferric hydroxide NH 4 C1 Ca (OH) 2 Mg(OH) 2 A1(OH) 3 Fe (OH) 3 75.8 215. 217. 297. 198 4. 3. Ferrous-ferric oxide .... Zinc oxide Fe 3 4 ZnO 265. 86 Cupric oxide ....... CuO 37 Mercuric oxide II gO 20.7 Silver oxide AffoO 6. Lead monoxide PbO 50, VI APPENDIX. sapiqding sajBqdsoqj sapixo sapjpoi saijBiuojqo APPENDIX. VU CO o g fe S rH l> c3 b- s rH rH O GO z- I S 0^ -* O5 ^ Oi co O5 CO O5 S s rH ^ t^ s rH 00 !>. s ^ rH S rH i 1C o5 o Oi ^ Oi -* O5 Th O5 CO O5 CO Oi CO Oi g g CM Oi 00 b- *O -HI CO rH 00 00 ^* CO rH ifi 00 CO o CM CM rH Oi O co * A i 35 S CO Ci CO 05 Oi CM Oi CM Oi S g CO rH CO 00 rH *o C5 rH fc. 00 *O o ^ fM 00 Oi rH co rH CO 1 g g CO Oi S GO T* Oi : rH g * s rH Oi rH 05 rH C5 rH O5 O Oi O Oi O Oi Oi GO Oi 00 Oi 00 H g 8 b- co O CO co O rH "ti rH os t>. $ 00 rH fe b- o s i rH Oi o o^ O O5 O OS Oi 00 Oi 00 Oi GO GO GO GO GO CO 00 b- 00 CM ^ rH 1 1 O5 l^ 00 ^* 00 rH So t^ O t^ CM b- Oi co GO CO r S i S Oi CO 28 00 00 88 fe b- 00 b- 00 e oi 00 s t^ rH OS 00 t^ 00 88 t^ iO GO CO CO (M 28 co Oi I>I 00 t^ GO co CO b^ GO s b^ GO 00 b- PH* M H H O o iQ rH & rH o rH rH 5 rH 1 o rH CN O s O co CM O ^ CM o M CM Vlll APPENDIX. Table IV. LIST OF THE METALS IN THE ORDER OF DECREASING SOLUTION TENSION. The Alkali Metals. The Alkaline-Earth Metals. Magnesium. Aluminum. Manganese. Zinc. Cadmium. Iron. Cobalt. Nickel. Tin. Lead. Hydrogen. Bismuth. Antimony. Copper. Arsenic. Mercury. Silver. Palladium. Platinum. Gold. The metals appearing first in the table can replace those that follow, in the solutions of their salts. INDEX.* TNumbers denote pages, or, if preceded by " Ex.," Laboratory Exercises in Tart II.] Acetic acid 95 Acetylene 207 Acid and normal salts, Ex. 22, Ex. 23 properties 93 radical 98 salts 100 salts, converted into normal salts 100 acetic 95 carbolic 210 carbonic 197 chloric 266 chlorous 266 chlorplatinic 405 lluosilicic 317 hydriodic, 261, 262, Ex. 44, Ex. 45 hydrobromic 259, Ex. 45 hydrochloric . 88, 89, 90, 91, Ex. 17 hydrocyanic 204 hydrofluoric 255 hydrosulphuric . 169, 170, Ex. 33 hypocblorous 265 hyponitrous 163 hypophosphorous 285 iodic 268 nitric ... 95, 149, Ex. 18, Ex. 26 nitric, fuming 154 nitrous 158 perchloric 267 periodic 268 phosphoric . . 286, 287, 288, Ex. 47 Acid, phosphorous ....... 286 silicic 318 stannic 404 sulphuric, 95, 174 /., Ex. 18, Ex.35 sulphurous 183 thiosulphuric 184 Acidity and basicity 102 Acids, action of, on bases .... 93 action of, on metals 93 action with oxides 99 most important property of . 96 nomenclature of 103 polysilicic 319 properties of Ex.18 see under specific names. strong and weak 330 Agent, drying, for hydrogen (calcium chloride) 12 drying, for hydrochloric acid (sulphuric acid) 89 drying, for ammonia (lime) . 137 Air, a physical mixture 119 determination of proportion of oxygen 117 liquefaction of 115 Alabaster 352 Alkalies, defined 97 Alkali metals 337 distinction between . . . Ex. 57 properties change in order of atomic masses 333 * This Complete Index is bound in the book with and without the Laboratory Exercises. ix INDEX. Alkaline earth metals 350 Allot ropism 167, 251, 272 Alum, crystallization of . . Ex.13 water of crystallization in, Ex. 10 Aluminates 380 Aluminum and its compounds, Ex. 69 bronze 368 carbonate and sulphide . . . 382 hydroxide .... Ex.69,d, g occurrence and preparation of 378 oxide and hydroxide 381 properties of 379 salts 382 uses of ...,,, 380 Alums 382, Ex.69, e,/ Amalgam, ammonium . 348, Ex. 57 Amalgamation process for silver 370 Amalgams 365 Ammonia, action of chlorine on 85, 146 a refrigerating agent 140 chemical properties of . ... 142 commercial sources of .... 138 composition of 145 existence of 136 from illuminating gas manu- facture 209 laboratory methods of prepa- ration 136, 137 liquefaction of . 140 physical properties of .... 139 preparation and properties of Ex. 25 process of making soda ... 341 synthesis of 145 Ammonium . . * 348 amalgam Ex. 57 compounds 142 compounds, dissociation of . 144 hydroxide 97 Amorphous carbon 187 artificial 188 Analogues of an element . . . 309 Analysis, volumetric, use of per- manganate in 394 Anhydride of acid defined ... 96 Aniline 157 Animal charcoal and bone- black 191 Annealing and tempering steel . 386 glass 321, Ex.2 Antidote for arsenic 294 Antimony . Ex. 49 black 300 chemical properties 298 compounds 299 physical properties 297 preparation 297 trisulphide 300 uses of 300 Aqua regia 82 a source of nascent chlorine . 154 Argon Ill family 314 Arsenates 296 Arsenic c , . . Ex. 48 acid 295 chloride 291 greens 295 Marsh's test for 292 occurrence and preparation of 290 properties of 290 trioxide 293 trisulphide 296 Arsenious acid 294 double nature of 294 Arsenite of sodium 294 Arsine 292 Asbestos 358 Ash 188 of land plants 201 of sea plants 201 Atomic hypothesis 224 masses, determination of, 233, 234 masses, exact, obtained by comparing with oxygen . . 235 masses, relative 232 mass methods, application of, 237-240 Atomic, or nascent, state . . 242,243 Atoms 224 INDEX. XI Atoms and molecules, distinc- tion between 226 number of, in molecules of elements 243 Atmosphere, carbon dioxide in, 112 character of the 109 water vapor in 113 weight and pressure of ... 114 Atmospheric dust 113 Avogadro's hypothesis .... 128 Baking: powders 195 Balanced, or equilibrium, equa- tion 263,265 Barium and strontium .... 356 Barium chloride, water of crys- tallization in Ex. 61 hydroxide 97 peroxide 356 peroxide, a source of oxygen, 27 Barometric reading, correction Of 126 Bases, nomenclature of 105 properties of Ex. 19 Basic salts 101 Basicity and acidity 102 Bead, borax Ex. 51 Bell-metal 368 Benzene 210 Beryllium, or glucinum .... 358 Bessemer converter, Fig. 63 . . 387 process for steel 387 Bicarbonates- 198 Bismuth 300, Ex. 50 properties of 301 salts 301 uses of 302 Blast-furnace, Fig. 62 384 Bleaching powder 355 with chlorine ... 86, 87, Ex. 16 with sulphur dioxide .... 182 Blowpipe, oxidation of tin, Ex. 75, h oxidizing and reducing flames, Ex. 63, b oxyhydrogen Fig. 6, 14 reduction of lead oxide . Ex. 74, 6 Blue prints ..,.,..,... 374 Blue vitriol, crystallization of, Ex. 13 water of crystallization in, Ex. 10 Boiling-point of water 48 Bone-black and animal charcoal, 191 Borax 323 and boric acid Ex.51 Borax bead Ex.51 Boric acid 323 Boron, occurrence and prepara- tion 321 properties of 322 Boyle's, or Mariotte's, law . . 121 Brass 368 Bromates 268 Bromine Ex. 43 oxyhydrogen compounds . . 267 preparation of 257 properties of 258 Bronze 368 Bunsen burner 217, Ex. 1 flame 217, Ex. 41 Burner, Bunsen Ex. 1 Burning. See Combustion. Cadmium 363 compounds Ex. 65 Calcium 350 carbide 207 carbonate 352 chloride 352 chloride, deliquescence of, Ex. 12 compounds Ex. 58 hydroxide 97, 351 oxide 350 phosphate 353 silicate 356 sulphate 352 sulphide 356 Calculation of quantities of fac- tors and products 20 Calomel 364 Candle flame, Figs. 51 and 52, 215, 216 " Carats fine " . 377 Carbon 186, Ex. 37 amorphous 187 amorphous, artificial .... 188 Xll INDEX. Carbon and hyd-regen compounds, 205 coke and gas 190 reduction by Ex. 40 Carbonates 197 identification of 198 natural 200 Carbon dioxide ... Ex. as, Ex. 39 chemical properties of . . . . 193 in the atmosphere 112 occurrence of 191 other sources of 194, 195 ph}-sical properties ot .... 192 preparation 192 relation of, to life 196 Carbon disulphide 174 Carbon monoxide, formation of, 201 laboratory method 202 properties of 202 reduces iron ore 3S5 Carbonic acid 197 Carbonic anhydride 197 Carbonization, natural; coal . . 188 Carborundum 318 Cast-iron 385,386 Caustic potash 345 Caustic soda 340 Celluloid 157 Cement and mortar 355 hydraulic 355 Portland . 355 Chalk 352 Chamber crystals ....... 176 Change, chemical Ex. 3 chemical, and energy changes, 219 Changes, quantitative cnaracter of chemical 19 Charcoal, animal, and bune-black, 191 properties of Ex. 37 reduction of lead oxme by, Ex. 40, a wood 189 Charles' law 123 Charring ....' 187 Chemical change Ex. 3 fundamental fact of 66 Chemical changes aim er.ergy changes 219 Chemical changes, quantitative character of 19 reaction defined 66 Chemistry and physics, relation between . 2 definition of s 1 importance of 1 Chile saltpeter 156,338,344 Chlorates and chloric acid . . 206 Chlorides 93 formed by dissolving metals in aqua regia 82 Chlorine, action of, on ammo- nia 85, 146 and hydrogen, union of ... 84 and turpentine 86 and water 85 chemical properties of .... 84 common method of prepara- tion 79 existence of 79 liquefaction of 83 other methods of prepara- tion 81, 82 preparation and properties of Ex. 16 uses of .' 86 water, preparation of Ex. 16, c, i Chlorine dioxide and tetrox- ide 264 Chlorine hydrate 83 Chlorine monoxide 264 Chlorous acid 2(56 Chlorplatinic acid 405 Chromates and dichromates, 398, 399 distinguished from chromic salts Ex.73,/ Chrome-alum 397 Chromic salts changed to chrom- ites Ex. 73, e distinguished from chromates, Ex.73,./ from dichromates or chrom- ates Ex. 73, h Chromite . 395 Chromites . ... 397 INDEX. Xlll Chromium 395 compounds Ex. 73 double nature of 397 oxides and hydroxides . ... 390 Chromous and chromic salts, 396 Clay 378, 3S3 Coal, natural carbonization ... 188 " Coal gas " 203 Coals and wood, composition of, 189 Cobalt 391 Cobalt and nickel Ex. 71 chloride and water . . . Ex. 71,/ Coin gold 376 Coin silver 372 Coke 210 Coke and gas carbon 190 Collodion 157 Couibining proportions .... 73 use of 74 Combustion, fixed amount of heat of 31 in air 15 in air; drafts 33 in chlorine 84 in oxygen . Ex. 7 oi'dinary 31 reversed 36 slow 31 spontaneous 32 supporting 15 temperature depends on rate of 31 Commercial iron 386 Comparison of the halogen acids Ex. 45 Compound, defined ....... 5 of two elements, how named, 74 Compounds, molecules of ... 227 of same two elements, how distinguished 75, 77 Conductivity, electric, of solu- tions 326 Constant proportions . . . Ex. 15 law of 67 Converter, Bessemer, Fig. 63 . . 387 Copper 366 alloys 368 Copper and its compounds . Ex. 67 compounds 368 etc., relation to alkali metals . 366 native .'.... 366 ores 307 properties and uses 367 sulphate use in making hydro- gen 11 Copper-plating 369 Coral 201 Correction of the barometric reading 126 Corrosive sublimate 365 Critical temperature defined . 116 Crucible process for steel . . . 387 Cryolite 378 Crystallization 63, Ex. 13 from fusion 64 water of Ex.10 Crystal-water in barium chlor- ide . Ex.61 in gypsum Ex.59 Cupric oxide 369, Ex. 67, c Cupric sulphate 3(59 Cupric sulphide 369 Cuprous oxide .... 368, Ex. 67, e Cyanide process of extracting gold 375 Cyanogen 203 Davy's safety lamp 35 Deacon's process for chlorine . 82 Decantation and filtration . . 64 Decay of wood 31 Decomposition, heat of .... 220 Decrepitation 54, Ex. 23 of salt 344 Definite proportions, law of, explained by atomic theory, 224 Deflagration 30 of potassium nitrate and sul- phur Ex. 55 Dehydrating or drying agents, 12, 56, 89, 137, 346 Deliquescence .... 56, 131, Ex. 12 Density, methods of vapor . . . 228 XIV INDEX. Developing, photographic ... 373 Dew-point 113 Diamond 186 Diatoms 315 Dichromates and chromates, 398, 390 Diffusion, explanation of . ... 129 of hydrogen 17 Dioxides and peroxides .... 277 Disodium hydrogen phosphate and magnesium salts, Kx . 62, d Dissociation by heat 325 iu solution, ionization .... 325 of ammonium compounds . . 144 of arsenic molecules 291 of hydi-iodic acid 262 of hydrobromic acid 260 of iodine 261 of nitrogen tetroxide 160 of steam 48 Distillation, apparatus for ... 46 defined 45 Distilling at reduced pressure, 275 Dolomite 201 Double cyanide of silver and potassium Ex.68,/ Drafts, combustion in 34 Drying agent 56 calcium chloride . . . 12,137,352 calcium oxide 56 potassium carbonate 346 sodium hydroxide ...... 137 sulphuric acid 89 Dulong and Petit's rule .... 235 Dust, atmospheric 113 Dynamite 157 Earthenware 383 Effervescence 04 effect of pressure on ... Ex. 39 Efflorescence 55, 131, Ex. 11 Electric conductivity of solu- tions 326 Electric current, to extract met- als from compounds .... 335 Electric furnace, Figs. 48 and 61, 207, 379 Electrolysis 329 of hydrochloric acid, Fig. 23, 83, 91 of sodium chloride 83 of water 19 preparation of aluminum by, 370 preparation of potassium by . 345 preparation of sodium by . . 339 preparation of sodium hy- droxide by 340 Electrolytes 19, 327 Electrotype plates 369 Element, analogues of an ... 309 Elements, abundance of .... 7 defined 5 heterologous 300 homologous 305 importance of 7 list of 6, Appendix molecules of 227 number of atoms in molecule of 24.J periodic table of 312,313 prediction of unknown . . . 311 properties of, determined . . 309 Energy changes accompany chemical changes 219 potential, of elements .... 219 Equation, equilibrium or bal- anced 263,265 ionic 3_'S volumetric meaning of . . . . 246 Equations and molecular formu- las 241 and symbols, quantitative meaning of 70 symbolic 68 the result of experiment . . . 70 Equilibrium equation of a solu- tion 327 in solutions of gases and sol- ids 131,132 or balanced equation . . 263, 265 unstable, of compounds with negative heat of formation, 221 Equivalent 237 of magnesium ....... Ex. (I INDEX. XV Equivalent of zinc . . . 237, Ex. 64 Etching of glass 256,321 Ethane 206 Ethylene 206 Eudiometer 40 Evaporation 131, Ex. 4 Explosion of hydrogen, velocity of 16 Factor denned 4 Factors and products . . 20, 69, 80 Family, argon . . . 314 calcium 359 halogen 268 nitrogen 303 Families, natural 305 Feldspar 378 Fermentation 194 Ferric chloride 390 Ferric hydroxide 389 Ferric oxide 389 Ferric sulphate 390 Ferrous ammonium sulphate, 390 Ferrous chloride 389 Ferrous-ferric oxide ...... 389 Ferrous hydroxide ....... 389 Ferrous sulphate 390 Fertilizers 354 Filtrate 64 Filtration Ex. 4 Fire-damp 205 Fixing, photographic 374 Flame, Bunsep, Fig. 54 217 candle, Figs. 51 and 52 . . 215, 216 colored by barium . . 357, Ex. 60 colored by calcium . . . . Ex.58 colored by potassium, 345, Ex. 55 colored by sodium . . 344, Ex.54 colored by strontium . 357, Ex. CO defined 35, 214 hydrogen, Fig. 7 14 non-luminous 216 oxidizing, with blowpipe, Ex. 63, 6 reducing, with blowpipe, Ex. 63, 6 Flames Ex. 41 dissection of, Fig. 55 218 luminosity of , 214 Flames, oxidizing and reduc- ing Ex.51 simple and complex 217 structure of 215 Flask, generating 10 Fluorine 254,255 Fluosilicic acid 317 Flux 254 used in blast-furnace .... 384 Fool's gold 389 Formula types based on valence, 248 Formulas and symbols 68 and symbols, how to rcpi'escnt multiples of 71 graphic, or structural . . 249, 25J how determined 240 molecular, and equations . . 241 Freezing-point of water .... 48 Fuming nitric acid 15t Furnace, blast, Fii,'. 62 384 electric, Figs. 48 and 61 . 207, 379 Hall's aluminum, Fi^. 61 . . 379 reverberatory, Fig. 60 .... 371 Gallium, properties predicted . 311 Galvanized iron 362 " Gaps " in the periodic arrange- ment 310 Gas, amount used 213 carbon 210 carbon and coke 190 collection of, over water, 10, Ex. 5 collection of, by upward dis- placement 11 collection of, by downward displacement 80 comparison of the two kinds of illuminating 212 permanent 116 relation of pressure to volume of 121 relation of temperature to vol- ume of 122 Gaseous substances, molecular masses of 227 Gases and vapors defined .... 121 XVI INDEX. Gases, diffusion of 120, 130 solution of, in liquids .... 131 Generating flask 10 Generator, Kipp's, Fig. 4 .... 12 Germanium, properties pre- dicted 311 German silver 36S Glass 320 color of 320 cut 321 etching of 250, 321 prcs^el 321 tubing, cutting and bending, Ex. 2 Glauber's salt 344 effloresce nee of Ex. 11 Gluciimtn 358 Glycerine 340 Gold 375 fineness, in carats 377 properties and uses of .... 376 Granite 378 " Granite ironware " 320 Grape-sugar and potassium permanganate . . . Ex.72, d Graphic, or structural, formulas, 249, 250 Graphite 187, Ex. 37 retorts used in distilling zinc, 371 Greens, arsenic 295 Green vitriol 390 Group, calcium 359 zinc . 361 Gun cotton 157 metal 368 powder 157 Gypsurn 352 water of crystallization in, Ex. 59 Haematite 384 Hall's process for making alu- minum 378 Halogens 254 Halogen acids, comparison of, Ex. 45 family 268 Halogen-oxygen compounds . 263 Halogen oxyhydrogen conn- pound* 264,265 Hardness, permanent 47 temporary 47 Heat, dissociation by 325 of formation and decomposi- tion 220,221 of formation evolved in stages, 221 of formation, positive and negative -221 Heating liquids in beakers, etc., Ex. 4 with Bunsen flame, best method qf Ex. 1 Helium Ill Heterologous elements .... 309 Homologous elements .... 305 Hydraulic cement 35.") mining of gold 375 Hydriodic acid . 261, Ex. 44, Ex. 45 powerf ul reducing agent . . . 2G3 properties of 262 Hydrobromic acid, . . . 259, Ex. 45 Hydrochloric acid, commercial method of preparation . *s> 89 common method of prepara- tion 88 electrolysis of 91 existence of 88 physical properties 90 preparation and properties, Ex. 17 synthesis of 91 volumeti ic composition of . . 91 Hydrocyanic acid 204 Hydrofluoric acid 255 Hydrogen and chlorine, union of 84 antimonide 29!) rsenide 292 methods of preparation ... I'.) occlusion of 16 peroxide 274, Ex. 46 peroxide, composition of . . 276 peroxide, test for 276 phosphide, Fig. 58 282 physical properties of .... 17 preparation and properties of Ex. 5 reducing power of ...... 10 INDEX. XV11 Hydrogen silicule 316 sulphide 169, Ex.33 sulphide, properties of .... 170 velocity of explosion of ... 16 Hydrogen-carbon compounds, 205 Hydrolysis 330,331 and mass action Ex.53 of aluminum carbonate and sulphide 382 of magnesium carbonate . . 358 of soap 341 Hydrosulphides 172 Hydrosulphuric acid 1G9 Hydroxides 52 action of metals upon .... 52 Hydroxyl 98 Hydroxylamine . . 152 Hygroscopic 56 "Hypo" 184 Hypobromites 267 Hypochlorites 265 Hypochlorous acid 265 Hyponitrous acid 163 Hypophosphorous acid .... 285 Hypothesis, Avogadro's .... 128 atomic 224 Iceland spar 201 Ignition temperature 32 Illuminating gas 208 comparison of the two kinds . 212 distillation of coal, Fig. 49 . . 208 water gas process, Fig. 50 . . 210 Ink, sympathetic 391 lodates 268 lodicacid 268 Iodine Ex. 44 occurrence and preparation . 260 Iodine oxyhydrogen com- pounds 268 Iodine pentoxide 264 properties of 261 Ionic equation 328 lonization 325, 327, Ex. 52 Ions .T. 7 . . 327 not atoms 330 Iron Ex. 70 Iron chlorides 389 commercial 386 compounds, comparison of the two classes Ex. 70, e galvanized 362 occurrence and metallurgy . 384 oxides and hydroxides .... 388 properties of 388 pyrites 169, 384, 389 sulphates 3!)0 sulphides 389 Isouierisml 250 Kaolin 319,383 Kindling temperature . 32, Ex. 8 Krypton 314 Lampblack 189 Latent lieat of water 48 Law, TSoyle's or Mariotte's . 121,122 Charles' 123 of conservation of matter . . 67 of constant proportions by weight 67 of definite proportions ex- plained by the atomic the- ory 224 of multiple proportions . . . 223 of multiple proportions, ex- planation of 225 of persistence of mass .... 67 Laws of osmotic pressure . . . 133 of simple and multiple vol- umes 215 Lead and its compounds . . Ex. 74 occurrence, preparation, and properties 400 oxides, nitrate and acetate . . 401 sulphate, chloride, and car- bonate 402 sulphide 402 uses of 401 Le Blanc process 89, 341 Life, relation of carbon dioxide to 196 Light, action on silver com- pounds 373, Ex. 68 XV111 INDEX. Lime 350 "air slaked" 351 slaking Ex.58 Limestone 200, 352 Liquefaction of ammonia . . . 140 Liquid air, properties of .... 117 Liquids, diffusion of .... 129,130 Lithium 338 Litmus solution, effect of char- coal on Ex. 37 Lodegtone 389 Luminous flames Ex. 41 Luster, metallic 334 Magnalium 381 Magnesite 201 effect of heat upon, Ex. 40, Ex. 62, e Magnetite 384, 389 Magnesium . 357 ammonium phosphate . . Ex.47 and its compounds . . . . Ex.62 carbonate 358 chloride 358 equivalent of Ex.6 oxide 357 silicide 317 sulphate 358 Manganatea 393 Manganese 392 compounds Ex. 72 dioxide and hydrogen per- oxide Ex. 46 dioxide a source of oxygen . 27 dioxide, action on potassium chlorate 26 dioxide, to liberate chlorine, Ex.16 dioxide, to prepare oxygen, Ex. 7 oxides 392 Manganites 395 Marble 201,352 result of heating Ex.58 Mariotte's, or Boyle'g, law, 121,122 Marl 352 Marsh gas 205 Marsh's test for antimony ... 299 for arsenic 292 Mass, persistence of 66 Mass action 331,332 and hydrolysis Ex.53 Matches 281 Matter, law of conservation of - 67 physical states of 123 Mercuric chloride 365 oxide 364 oxide, decomposition of, 27, Ex. 3 sulphide 365 Mercurous chloride 364 oxide 364 Mercury 363 and its compounds .... Ex. 66 compounds 364 Metal, Rose's 302 type 297 Wood's 302 Metallic luster 334 Metalloids 331 Metals, alkali 337 alkali, distinction between, Ex. 57 alkaline-earth 350 and non-metals 334 extraction of, from ores . 335, 385 occurrence of 334 properties of 335, 336 " Metaphosphate head " . '. . 288 Methane 205 Mica 378 Mineralg 334 Mirrors 373 Molecular formulas and equa- tions ... 241 mass, boiling-point and freez- ing-point methods 230 mass, exact methods of obtain- ing 231 mass of oxygen, 32, reason for . 244 mass, other methods of deter- mining 229 mass, oxygen the standard of, 228 mass, vapor density methods for 228,2-29 masses of gaseous substances, 227 theory 127 INDEX. XIX Molecules and atoms, distinction between 226 of elements and of compounds 227 of elements, number of atoms in 243 Mordant, aluminum compounds, 382 Mortar, action of acid on old, Ex. 58 anil cement 355 Multiple proportions, explana- tion of law of 225 law of 223,225 Multiples of symbols and for- mulas, how represented ... 71 Naming a compound of two ele- ments 74 different compounds of same two elements 75-77 Naphthalene 210 Nascent, or atomic, state . . 242,243 Natural families 305 Natural family of elements . . 270 Neon 314 Neutralization 98, Ex. 21 explanation of 327, 328 Neutralize 96 Nickel 391 and cobalt Ex.71 Nitrate, manufacture of potas- sium ..." 156 Nitrate beds 156 Nitrates and nitric acid, uses of 157 Nitrates, formation of, in nature, 155 Nitre . . . . T 346 Nitric acid 95,149 action on metals 152 and nitrates, uses of 157 commercial preparation . . . 149 explanation of oxidation by . 154 faming b 154 laboratory preparation of . . 150 oxidation by 154 powerful oxidizing agent . . 152 preparation Ex. 26 preparation compared with that of hydrochloric acid . . 150 properties of . . 151, Ex. 18, Ex. 26 Nitric acid, oxidizing agent, Ex. 34, e; Ex. 70, d; Ex. 75, 9 reduced by nascent hydrogen, 152 Nitric anhydride 158 Nitric oxide 161, Ex.29 Nitrides 109 Nitrites 159, Ex. 27 Nitrobenzene 157 Nitrogen dioxide and tetroxide, 160, Ex. 28 existence of . . . , 107 family 303,304 from ammonium nitrite . . . 108 of the air, a mixture 109 oxides 155 pentoxide 158 preparation and properties, E*. 24 preparation of, from air, Figs. 25 and 26 107 properties of 108 trioxide 109 Nitrous acid 158 anhydride 158 oxide 102, Ex. 30 Nomenclature of acids .. ... 103 of bases 105 of salts 104 Non-luminous flames . 21(5, Ex.41 Normal salts 99 and acid salts . . Ex. 22, Ex. 23 Occlusion of hydrogen .... 16 Open hearth process for steel, 387 Ores 334 extraction of metals from . . 335 Osmotic pressure . 132 laws of 133 Oxidation by chromates and di- chromates 399 by nitric acid . . . 154, Ex. 75, e by nitric acid, explanation of, 154 by permanganate 394 and reduction 30 Oxides 29 of phosphorus 284 with acids, action of ..... 99 A INDEX. Oxidizing agent ........ 30 bromine ........... 258 chlorine ........ Ex. 73, g chromates and dichromates . 399 hydrogen peroxide ...... 275 nitric acid ..... 154, Ex.75, e ozone ............. 273 permanganate ..... ... 394 Oxidizing flame ...... Ex.51 Oxygen acids of phosphorus . . 285 chemical properties of ... 28 determination of the propor- tion of, in air, Figs. 30 and 31 ............ 117,118 physical properties of .... 27 preparation and properties, Ex. 7 preparation of .... 24, 25, 26, 27 reason for choosing molecular mass of, 32 ... ...... 244 standard of atomic mass ... 232 standard of molecular mass . 228 weight of one liter .... Ex. 42 Ozone .............. 273 Palladium, occlusion of hydro- gen by ........... 17 Paris green ......... . . 295 Parke's process of desilverizing lead ............ 371 Peat ....... . ....... 188 Perchlorates .......... 267 Perchloric acid ..... .... 267 Periodic acid .......... 268 Periodic arrangement . . . 306JT. conclusion .......... 312 gaps in ............ 310 regularitiein ........ 308 Periodic law .......... 308 Permanent gases ...... *. . 116 Permanganate, oxidation by . 394 Peroxides and dioxides .... 277 Peroxide of hydrogen . . . Ex. 46 Phenol ............. 210 Phenomena, chemical and phys- ical ............. 3 Phosphates, necessary for plants, 354 uses of . ........... 289 Phosphide of hydrogen, Fig. 58 ............... 282 Phosphides ........... 284 Phosphine, Fig. 58 ....... 282 Phosphonium salts ...... 283 Phosphorescence ....... 168 Phosphoric acids ....... 28(i preparation of ........ 287 salts of ............ 288 Phosphorous acid ....... 286 Phosphorus ........ , . 279 and phosphoric acid . . . Ex.47 molecular mass of ...... 281 oxides ............ 284 properties of ......... 279 red .............. 280 Photography ........ . . 373 Physical phenomena ..... 3 Physics, relation between Chem- istry and ........... 2 Pig-iron ...... , ...... 38;i Placer-mining of gold ..... 375 Plaster of Paris ........ 352 making casts with .... Ex. 58 Platinum, occlusion of hydrogen by ............ 16, 17 occurrence and preparation . 404 properties and uses ..... 40.1 Plumbites and plumbates . . 402 Pneumatic trough !...... 10 Polysilicic acids ........ 319 Porcelain-lined ware ..... 320 Porcelain, stoneware, etc. . . 383 Portland cement ...*.... 355 Potash ........... 346, 347 Potassium ......... , . 345 bromide ........... 347 carbonate ........ 201, 340 chlorate ........... 346 chloride, solubility of . . Ex. 5<; compounds ........ Ex. 55 ferricyanide ......... 391 ferrocyanide * ....... 390 hydrogen tartrate . Ex. 55, Ex. 57 hydroxide ......... 97, 345 hydroxide, deliquescence of, Ex. 12 INDEX. XXI Potassium iodide 347 nitrate 346 nitrate, manufacture of, 156, Ex. 55 permanganate 393 silicate 319 Powders, baking 195 Precipitant 63 Precipitate 63 Precipitation Ex. 14 and crystallization 63 Prediction of unknown ele- ments 311 Pressure and temperature, re- duction to standard .... 124 osmotic 132 relation of, to volume of a gas, 121 Products and factors 69 calculation of quantities of factors and 20 Properties, chemical and physi- cal 4 effect of structure on .... 277 of acids Ex. 18 of bases Ex. 19 of elements determined . . . 309 of salts Ex. 20 Proportions, combining .... 73 constant 67, Ex. 15 law of multiple 223, 225 use of combining 73 Prussian blue .... 391, Ex. 70, d Pyrites 174, 389 Pyrolusite . _." 27, 392 Quantitative meaning of sym- bols and equations .... 70 Quicklime 350 Radical, acid 98 metallic 143 Reaction defined 4, 66 Reactions, classification of ... 66 Reagent defined 4 Red phosphorus 280 precipitate 364, Ex. 3 Reducing agent, hydriodic acid, 263 hydrogen peroxide 276 Reducing agent, nascent hydro- gen 152 Reducing flame . . Ex. 51, Ex. 63 Reduction 30 by carbon Ex. 40 by hydrogen 42 of potassium nitrate by lead, Ex. 27 Refrigerating agent, ammonia as a 140 Relation of volume of a gas to pressure 121 Reverberatory furnace .... 371 " Reversion to type " 307 Rose's metal 302 Safety lamp 35 .matches 281 Salt 343 acid 100 defined 99 "Epsom" 358 normal 99 Saltpeter 346 Salts, basio , . 101 nomenclature of 104 normal and acid . Ex. 22, Ex. 23 properties of Ex. 20 Saponify 340 Saturated solution 62 Scandium, properties predicted, 311 Scheele's green 295 Schweinfurt's green 295 Sea-water . 45 Semipermeable cell 133 Shells 201 Siderite, or spathic iro ore . . 384 Silicates 319 Silicic aoid 318 anhydride , .... 318 Silicon carbide , . . 318 compounds 316 dioxide 318 hydride 316 occurrence and preparation . 315 tetrachloride 317 tetrafluoride . .... 317 XX11 INDEX. Silver 370 and its compounds .... Ex. 68 bromide . . . 373 chloride 373 " coin " 372 compounds . . . 373 extraction of 370 iodide 373 nitrate 373 -plating 372 properties and uses *372 "sterling" 372 Slag 254 Slaking of lime 351, Ex. 58 Smelting of silver 370 Soap 340 hydrolyzed In solution .... 341 Soda-ash 342 Soda, crystallized 342 Sodium 338 action upon water . . . 49, Ex. 9 amalgam 92, 339 and water, quantitative study of the reaction 52 arsenite 294, Ex. 48 bicarbonate . . 343, Ex. 39, Ex. 54 carbonate 201, 341, Ex.11 chloride 343 compounds Ex.54 dichromate 399 hydroxide 90,340 nitrate 156,344 oxides 339 phosphate 343 preparation and properties, 49, 339, Ex. 9 sulphate 344, Ex. 11 sulphite Ex.34 zincate Ex. 63, d, e " Softening " water with goda, 353 Solids, diffusion of 129 solution in liquids 131 Solubility ...... 60 of potassium chloride . . Ex. 56 Soluble and insoluble substan- ces 61 Solution . . x. 4 Solution, boiling point of .... 59 character of ......... 58 composition of a 327 effect of temperature on, 60, Ex. 13 equilibrium equation of ... 327 freezing-point of 59 of gases in liquids 131 of solids in liquids 131 of starch Ex.44 saturated 62 supersaturated 62 temperature changes during , 59 Solutions, electric conductivity of 326 Solvay, or ammonia, process . . 341 Solvent 58 Specific heat defined . ( . . . . 285 of water 47 table Appendix Spelter 301 Stalactites, formation of, Fig. 46, 200 Standard temperature and pressure, reduction to . . . 124 Stannic acid 404 and staunous compounds, 404, Ex. 75 Starch solution Ex.44 States of matter, physical . . . 128 Steam and its dissociation ... 48 volumetric composition of . . 41 Steel, properties and manufac- ture of 3S6, 387 Sterling silver 372 Stibine 299 Stoneware, porcelain, etc. . . 383 Strontium and barium, 356, Ex. 60 Structural, or graphic, formulas, 249, 250 Structure, effect of, on properties, 277 Sublimation . 144 Substance, compound and ele- mentary . . . 5 Sulphantimonites 300 Sulpharsenious acid 296 Sulphates 179, Ex. 36 Sulphides, formation of .... 171 precipitation of 172 INDEX. XX111 Sulphites 183 Sulphostannates 404 Sulphur, chemical properties of, 167, Ex. 32 compounds of 1(>9 dioxide 181, Ex. 34 dioxide, properties of .... 182 occurrence and preparation . 165 physical properties of, 165, 166, Ex. 31 trioxide 180 uses of 168 water 44 Sulphuric acid 95 and sulphates, test for .... 180 diluting 88, 89, Ex. 17 hydrates of 177 manufacture of 174 properties of . 177, Ex. 18, Ex, 35 purification of . 176 reduction of 177 uses of 179 Sulphuric anhydride 180 Sulphurous acid . 183 Sulphurous anhydride .... 183 Supersaturated solutions ... 62 Suspension Ex. 4 Symhol defined 5 Symbolic equations 68 Symbols, advantages of use of . 69 and equations, quantitative meaning of 70 and formulas 68 and formulas, how to repre- sent multiples of 71 Sympathetic ink 391 System, periodic 300 JT. Table of atomic masses . Appendix of specific heats . . . Appendix the periodic 312, 313 of metric equivalents, etc., Appendix of vapor tension of water, Appendix Tartar emetic 300, Ex. 49 Tart aric acid 96 Temperature and pressure, re- duction to standard .... 124 critical 116 effect of, upon solution, 60, Ex. 13 kindling 32 of ignition 32 relation of volume of a gas to, 122 . Tern pering and annealing steel, 386 " Test " reactions are ionic . . . 328 Theory, atomic 224 ionization 325 molecular 127 Thiosulphates 184 Tin and its compounds . 403, Ex. 75 occurrence and preparation of 402 properties and uses 403 Tincture Ex. 44 Toluene 210 Toning, photographic 374 Transition elements .... 295, 299 Transpiration . . < 18 Trough, pneumatic 10 Tube, safety, or thistle 10 Tubing, cutting and bending glass Ex. 2 Turnbull's blue Ex. 70 Type, reversion to 307 Type-metal 297 Ultramarine 383 Valence 241 different formula types based on 248 of members of argon family, 314 Vapor density methods . . 228, 229 Vapors and gases defined ... 121 Vein-mining of gold 375 Victor Meyer's method for va- por density 229 Volume of a gas, relation of pres sure to 121 of a gas, relation of tempera- tureto 122 Volumes, laws of simple and multiple 245 XXIV INDEX. Volumetric analysis, use of per- manganate 394 composition of steam .... 39 composition of hydrochloric acid 91 composition of ammonia, 145,146 meaning of an equation ... 246 Water, action of sodium upon, 49, 50, Ex.9 and chlorine, action of .... 85 and potassium 51 as integral part of a substance, 55 distilled 47 electrolysis of ....... 19, 38 formation of 15 hard and soft 47 in combination 53 mechanically enclosed .... 53 natural, and its impurities . . 43 nature of . . 38 of crystallization ... 54, Ex. 10 of crystallization in barium chloride Ex. 61 of crystallization in gypsum, Ex.59 properties of 47 purification of 45 Water, sea ,.... 45 " softened " with soda .... 353 sulphur 44 synthesis of, by weight . . . 42 synthesis of, by volume ... 39 vapor in the atmosphere ... 113 -gas process 210 -glass 319 Weight, constant proportions by 67 of one liter of oxygen . . Ex. 42 White-lead 402 Wood charcoal 189 Wood's metal 302 Wrought-iron 386 Xenon 314 Zinc 361, Ex.63 chloride 363 compounds 362, Ex. 63 dust 361 equivalent of Ex. 64 group 361 oxide 362 sulphate and sulphide .... 363 used in desilverizing lead . . 371 uses of 362 THIS BOOK IS DUE ON THE LAST DATE STAMPED BELOW AN INITIAL FINE OF 25 CENTS WILL BE ASSESSED FOR FAILURE TO RETURN THIS BOOK ON THE DATE DUE. THE PENALTY WILL INCREASE TO SO CENTS ON THE FOURTH DAY AND TO $1.OO ON THE SEVENTH DAY OVERDUE. Jtinu ^ n V -9 193S f\^f c\c\ < /s. 22 !& if! *, .. WAY $1 <| Jg 4n HjF OCT 9 19' 6 LD21-100m-7,'39(402s) f THE UNIVERSITY OF CALIFORNIA LIBRARY