SOLUBILITIES OP BASES AND SALTS IN WATER AT 18. K Na Li Ag Tl Ba \ Sr Ca Mg Zn Ptt Cl 82.95 3.9 35.86 6.42 77.79 13.3 0.0,13 0.0 5 9 0.3 0.013 37.24 1.7 51.09 3.0 73.19 6.4 55.81* 6.1 203.9 9.2" ss 1.49 0.06 Br 66.86 4.6 88.76 6.9 168.7 12.6 0.0 4 1 ' 0.0 8 6 0.042 0.0,15 103.6 2.9 96.52 3.4 143.3 6.2 ^ X*8fi. 103.1 4.6 478.2 9.8 0.598 0.02 I 137.6 6.0. 177J 8.1 161.5 8.5 0.0.35 0.0 7 1 0.006 0.0,17 201.4 3.8 169.2 3.9 200 4.8 148.2 4.1 419 6.9 0.08 0.0,2 P 92.66 12.4 4.44 1.06 0.27 0.11 195.4 , 13.5 ' 72.05 3 0.16 0.0,92 0.012 0.001 0.0016 0.0,2 0.0087 0.0,14 0.005 0.0,5 6.06 0.002 HO. 30.34 2.6 83.97 7.4 71.43 7.3 213.4 8.4 ' 8.91 0.35 8.74 0.33 66.27 2.7 121.8 6.2 74.31 4.0 117.8 4<7 61.66 1.4 CIO. 6.6 0.52 97.16 6.4 313.4 15.3 12.25 0.6 3.69 0.13 35.42 1.1 174.9 4.6 179.3 5.3 126.4 4.7 183.9 5.3 150.6 3.16 BrO, 6.38 0.38 36.67 2.2 152.5 8.20 0.59 0.025 0.30 0.009 0.8 0.02 30.0 0.9 85.17 2.3 42.86 1.5 58.43 1.8 1.3 0.03 10. 7.62 0.35 8.33 0.4 80.43 3.84 0.004 0.0,14 0.059 0.0,16 0.05 0.001 0.25 0.0,57 0.25 0.007 6.87 0.26 0.83 0.02 0.002 0.0 4 3 OH> 142.9 18 116.4 21. 12.04 6.0 0.01 0.001 40.04 1.76 3.7 0.22 i 0.77 0.063 0.17 0,02 0.001 0.0,2 0.0,5 0.0.5 0.01 0.0,4: S0 4 - 11.11 0.62 16.83 1.15 35.64 2.8 0.55 0.020 4.74 0.09 0.0.23 0.0 4 10 0.011 0.0,6 0.20 0.015 35.43 2.8 53.12 3.1 0.0041 0.0,1? GrO 4 - 63.1 2.7 61.21 3.30 111.6 6.5 0.0025 0.0,16 0.006 0.0,1 0.0 3 35 0.0 4 14 0.12 0.006 0.4 0.03 73.0 4.3 .* I I 0.0 4 2 o.oj? 0,0, 30.27 1.6 3.34 0.24 7.22 0.69 0.0034 0.0,17 1.48 * 0.030 0.0085 0.0,38 0.0046 0.0 3 26 0.0 3 55 0.0 4 43 0.03 0.0027 0.0 S 64 0.0 4 4 0.0,16 0.0,54 CO,- 108.0 5.9 19.39 1.8 1.3 0.17 0.003 0.0,1 4.95 0.10 0.0023 0.0,11 0.0011 0.0 4 7 0*0013 0.0,13 0.1 0.01 0.004? 0.0,3? 0.0,1 0.0 4 3 The upper number in each square gives the number of grams of the anhydrous salt held in solution by 100 c.c. of water. The lower number i the molar solubility, i.e., the number of moles contained in one liter of the saturated solution. The numbers for small solubilities have been abbreviated Thus 0.0.4 - 0.0&06004. For some other solubilities, see p. 109 A tbe same autbot EXPERIMENTAL INORGANIC CHEMISTRY Fifth Edition, Revised INORGANIC CHEMISTRY Third Edition (/(?//), Rewritten 925 + xvi pp. With 166 Figures Octavo, A briefer Text on the same plan GENERAL CHEMISTRY FOR COLLEGES Second Edition, Revised 662 + xi pp. With 138 Figures LABORATORY OUTLINE OF COLLEGE CHEMISTRY 206 + v pp. With 30 Figures A Text-book of ELEMENTARY CHEMISTRY 439 + viii pp. With 98 Figures and 6 Plates A LABORATORY OUTLINE OF ELEMENTARY CHEMISTRY 137 PP- With 18 Figures NEW YORK, THE CENTURY CO. LONDON, GEO. BELL AND SONS M. W. LOMONOSSOV Russian author, statesman, and chemist, 1711-1765 INTERMEDIATE TEXT BOOK OF CHEMISTRY BY ALEXANDER SMITH H_BAX> OF DEPAKTJLE2JT OF CHEMISTBY, COLUMBIA UNIVEBSITY NEW YORK THE CENTURY CO 1920 : COPYRIGHT, 1919, BY THE CENTURY CO. Printed, April, 1919 Reprinted, August, 1919 Reprinted, September, 1920 PREFACE THIS text book of chemistry is described as Intermediate, because it is longer than the Elementary Chemistry, and shorter than the College Chemistry. It includes a brief outline of Agri- cultural Chemistry. It contains, also, a discussion of Foods and their heating values, an explanation of the methods for softening water, and other practical applications of chemistry. Tlpse who use the book need not necessarily teach all that it contains. Selec- tions can be made, according to the purpose of the course. A Laboratory Outline will be prepared for use along with the text. ALEXANDER SMITH New York, January 24, 1919. 435146 111 CONTENTS CHAPTER PAGR I. SUBSTANCES AND PROPERTIES 1 II. CHEMICAL CHANGE AND THE METHODS OF STUDYING IT 12 III. AIR AND OXYGEN. 26 IV. THE MEASUREMENT OF GASES. THE MOLECULAR HY- POTHESIS 44 V. HYDROGEN 50 VI. WATER 59 VII. CHEMICAL UNITS OF WEIGHT. FORMULA 73 VIII. MAKING OF FORMULAE AND EQUATIONS 98 IX. SOLUTION 107 X. HYDROCHLORIC ACID. CALCULATIONS 119 XI. CHLORINE. CALCULATIONS 133 XII. ENERGY AND CHEMICAL CHANGE 147 XIII. SODIUM AND SODIUM HYDROXIDE 158 XIV. ACIDS, BASES, AND SALTS 163 XV. VALENCE 181 XVI. CARBON. THERMOCHEMISTRY 188 XVII. THE OXIDES OF CARBON 198 XVIII. COMPOUNDS RELATED TO PETROLEUM. FLAME 208 XIX. STARCH AND COMPOUNDS THAT CAN BE MADE THEREFROM 219 XX. CHEMICAL EQUILIBRIUM 228 XXI. NITROGEN. THE ATMOSPHERE 245 XXII. AMMONIA 256 XXIII. SULPHUR AND HYDROGEN SULPHIDE 263 XXIV. OXIDES AND OXACIDS OF SULPHUR, PERIODIC SYSTEM. . . . 273 XXV. OXIDIZING SUBSTANCES 300 XXVI. NITRIC ACID 313 XXVII. THE HALOGEN FAMILY 323 XXVIII. PHOSPHORUS, ARSENIC, ANTIMONY, BISMUTH 335 XXIX. SILICON BORON 345 v VI CONTENTS CHAPTER PAGE XXX. COMPOUNDS OF SODIUM AND POTASSIUM 351 XXXI. THE RECOGNITION OF SUBSTANCES, I. A REVIEW OF THE NON-METALLIC ELEMENTS 362 XXXIi. FATS, SOAPS AND RELATED COMPOUNDS 371 XXXIII. EXPLOSIVES AND PLASTICS 382 XXXIV. CALCIUM AND ITS COMPOUNDS 387 XXXV. MAGNESIUM AND ZINC 415 XXXVI. ALUMINIUM 422 XXXVII. PLANTS, FUELS AND FOOD. AGRICULTURE 432 XXXVIII. IRON, COBALT AND NICKEL. 455 XXXIX. LEAD AND TIN 472 XL. COPPER AND MERCURY ; 480 XLI. SILVER, GOLD, PLATINUM 488 XLII. MANGANESE, CHROMIUM AND RADIUM 496 XLIII. RECOGNITION OF SUBSTANCES, II. METALLIC ELEMENTS. 505 APPENDIX. . 511 INTERMEDIATE TEXT BOOK OF CHEMISTRY CHAPTER I SUBSTANCES AND PROPERTIES WHEN exact information in regard to any sort of material is required, we hand the material to a chemist. To learn something about the nature of chemistry, let us watch the chemist at work: Properties. Suppose that the material is a piece of cloth, and we desire to know whether it is all wool, or partly cotton. The chemist places a piece of the cloth in a test-tube (Fig. 1), and pours in an amount of lye sufficient to cover it. Lye is a solution in water (98 parts) of a white solid, named by chemists, commonly, "caustic soda" and, more formally, sodium hy- droxide (2 parts). The contents of the test-tube are then heated over a flame and are kept at the boiling point for ten minutes. If the cloth dis- solves entirely, leaving a liquid, clear like water, then it was composed of nothing but wool. The chemist draws this conclusion because wool, al- though not affected by the boiling water, has the property of turning into a soluble substance when caustic soda is heated with it. If, on the other hand, the piece of cloth becomes thinner, obviously losing a part of its material, but leaving a part undissolved, then it contained cotton. This conclusion depends on the fact that cotton has the property of not being dissolved by caustic soda solution. 1 FIG. 1. TEXT BOOK OF CHEMISTRY The same conclusion could- have been reached in other ways. For example, some threads could have been taken from the edges of the sample and placed under a microscope. In this case, they would have been seen to be made up of long, hollow fibers or tubes (Fig. 2). But the cotton fibers (A) are smooth on the surface, while the woolen fibers (B) are covered with scales. By this difference in properties, the presence of both kinds, or of only one of the kinds, could quickly be found out. Still again, the chemist knows that wool will "take" almost any dye, while cotton remains uncolored by the greater number of dyes. He could, therefore, boil a piece of the cloth with a solution of a soluble dye such as red ink (eosin), or the red dye Ponceau 6 R for a 2\ JiSSl few minutes, and then wash the sample thoroughly in clean water. On examining the cloth with a microscope he could then observe F IG . 2. whether any fibers were still white. Here wool has the prop- erty of uniting with the dye, while cotton has not. The last plan could be used only with white goods, or goods not already strongly dyed. i The first method is the one which the chemist would probably employ in practice, because' by its means he can ascertain and report, not only the presence of cotton, but the proportion of cotton present. To do this, he weighs the dry piece of .cloth before placing it in the test-tube. Then, after the boiling with caustic soda solution, he washes what is left of the sample very thoroughly in running water, dries it, and weighs it again. The weight of this "residue" is that of the cotton. The difference between this and the original weight is the weight of the wool. The chemist is then able to state the percentages by weight of cotton and of wool in the original material. SUBSTANCES AND PROPERTIES 3 Substances. Upon considering these operations, we can discover a plan which the chemist has devised and employs in his work. Different samples of cloth are of different colors and appearance, even when they contain the same proportions of cotton and wool. A general inspection is, therefore, of little value. But in certain respects all samples of wool are alike, such as in dissolving in caustic soda, in "taking" certain dyes and in possessing a scaly surface. Those respects in which all samples of wool are alike are called the properties of wool. Simi- larly all samples of cotton are alike in certain respects, which are called the properties of cotton. And the properties of cotton and of wool are different, many of them very different, indeed. For the purpose of stating what he means, the chemist calls a kind of material, all specimens of which possess a certain set of properties, a substance. Wool is one substance,* and cotton another substance. Every part of a specimen of a substance has the same properties as any other part. If any portions can be found which have different properties, these are portions of another substance, accidentally or intentionally mixed with the first. The chemist calls the foreign matter an impurity, and the specimen an impure sample of the substance of which it is mainly composed. A substance, then, is a species or kind of matter, and all specimens of it show the same set of properties. The substance is recognized by its properties. A property such as we have in mind is more definitely called a specific property, or property common to all the members of a species of matter. A specific property is a property shown by all the specimens of a given substance. The plan which the chemist uses is that of ascertaining whether a given material is made up of one, or of more than one substance. He then describes it by naming the substances he finds in it. * In point of fact, wool contains several substances, but they are all alike in respect to the three properties mentioned above. They differ slightly in respect to other properties, and so can be distinguished from one another. 4 TEXT BOOK OF CHEMISTRY Another Illustration. This view-point is peculiar to the chemist. Each art or science has its own view-point its own way of thinking about a given object. By the geologist, a piece of granite is at once thought of as belonging to the older rock of the earth, and the geologist considers when it was formed and how it was formed (namely, by solidification of a melted material). To the builder, it is a very hard stone, expensive to cut and polish, but very ornamental, and very durable. How and when it was formed does not make any difference to the builder. To the chemist, as a chemist, on the other hand, the expense of cutting granite, and its ornamental character, are of no interest. Instead, the chemist notices at once that it is spotted, and, upon examining it closely, he observes that it appears to be a mixture. He breaks up a portion, and studies the properties of the fragments. Some are transparent like glass, are very hard, and in fact are in all respects like quartz or rock crystal (Fig. 3) . All specimens of rock crystal, from what- ever source, are alike in properties, FlG 3 FlQ 4 and quartz is, therefore, a distinct substance or species of matter. Again, certain of the particles in granite are dark, and with a penknife can be easily split into transparent leaves, thinner than paper. These fragments are in all respects like mica (sheets of which, under the name of "isinglass," are used to close the windows of stoves), which is another substance (Fig. 4) well-known to the chemist. Still others of the fragments are less transparent than the quartz and less hard. They can be split, but with much greater difficulty than the mica. They are crystals,* oblong in shape. These are pieces of a third substance, felspar (Fig. 5). All the particles in the granite belong to one or other of these three kinds. The chemist, then, studies the specific * Crystals are natural forms of a geometrical nature, assumed by solid substances (p. 111). SUBSTANCES AND PROPERTIES FIG. 5. FIG. 6. properties, such as the hardnesses and the crystalline forms of various parts of the specimen and seeks to state or describe the nature of the specimen in terms of the substances he finds in it. When flour is examined by the chemist, it appears to the eye to be all alike. Under the microscope, even, all he can learn is that it consists largely of grains, which have the characteristic appear- ance (first property) of grains of starch (see Fig. 66, p. 219). He places some flour on a square piece of cheesecloth and encloses it by tying with a thread (Fig. 6). On kneading the little bag in a vessel of water, the water be- comes milky. When the milky water stands, the white material settles to the bottom, the water can be poured off, and the deposit can be dried. This white sub- stance, when boiled with water, gives an almost clear liquid which jellies on cooling. This is another property of starch. A little tincture of iodine (solution of iodine in alcohol), dropped on a part of the starch, causes the latter to turn blue. This is a very characteristic property of (and therefore test for) starch. When the bag of flour is kneaded persistently in water which is frequently changed, the material finally ceases to render the water milky. The starch has all been washed out. When the bag is now opened, a sticky material is found in it. This is called gluten. The chemist therefore finds that the flour contains starch and gluten. He learns this by separating the components. The Law of Component Substances. That all materials can be described in terms of certain component substances, each of which has a definite set of specific properties, is the first and simplest law of chemistry, and at the same time the most fundamental. It was, perhaps, first clearly stated by Lomonossov (1742), 6 TEXT BOOK OF CHEMISTRY Mixtures. When a material contains more than one sub- stance, it is called a mixture. It will be seen that the chief fact about a mixture is that each of the component substances pos- sesses, while mixed with the others, the same properties which it has when alone. A component of a mixture does not alter any of the properties of any other component. Mixtures and Impurities. A material containing more than one component substance is called a mixture. The charac- teristic of a mixture is that each of the component substances, although mixed with the others, possesses exactly the same properties as if it were present alone. No'one of the components affects any other component, or alters any of its properties. Granite and flour are typical mixtures. When a specimen is composed mainly of one substance, and contains only minute amounts of one or more other substances, it is frequently spoken of as a specimen of the main substance containing certain specified substances as impurities. To be called an impurity, the foreign matter need not be dirty or offensive. Thus, common salt usually contains a little magnesium chloride, a white crystalline solid, as an impurity, and it is this impurity which becomes damp in wet weather. Again, compounds of lime and magnesium are common impurities in drinking water. Bodies or Specimens. It will be seen that substance is a general term, like the word "dog," covering the whole species. The substance iron includes all the iron in the Universe. When we refer to a particular piece of iron, we call it a body or a speci- men. If the body is homogeneous (all parts alike), it may be made of a single substance. If it is heterogeneous (differing in different parts) it is a specimen of a mixture like granite. Component. The ingredients of a mixture are called the components (Latin, put with), because they are simply placed together, without change, and can be separated without change. SUBSTANCES AND PROPERTIES 7 Other Specific Properties. Beside the specific properties which happen to have occurred in these illustrations, there are others which are constantly found useful by the chemist. Thus, in the case of solids, besides the hardness and crystalline form, he gives special attention to the temperature at which the substance melts (the melting-point), the specific gravity or density (weight of 1 c.c., see Appendix I), the color and the solubility or non-solu- bility in water. In the case of liquids, the temperature at which the liquid boils under atmospheric pressure (the boiling-point), the specific gravity, the mobility, the odor and the color (if any) are never exactly the same for any two liquids. We shall learn more about these properties as occasions for using them arise. Attributes and Conditions. It should be noted that the mass of a specimen is not a property, or, at least, not a specific property. Each specimen of a given substance has a different mass, although it has identical specific properties. The mass, although unchangeable, is a quality of a particular specimen only. So is it with the volume and the dimensions of a specimen. These are attributes of a specimen. Again the temperature is not a property. There is no particular temperature peculiar to quartz. Even the very same specimen may be at different temperatures at different times. Temper- ature and pressure are variable and are called conditions. The temperature or the pressure of a specimen may be changed at will. But the specific properties of a substance cannot be changed, so long as we have to do with the same substance. Law of Chemical Change. When the wool was boiled with caustic soda solution, it was in some way acted upon by the caustic soda. It became soluble, and disappeared into the liquid. Wool will not dissolve in boiling water. But in the former operation its material acquired at least one new property, namely that of being soluble. Since we have defined a substance 8 TEXT BOOK OF CHEMISTRY as "a species of matter, with a definite and constant set of proper- ties," we are compelled to decide that when a piece of material changes its properties, it has, in doing so, become a new substance. This experiment, then, calls our attention to the second of the fundamental laws of chemistry, namely that material forming one or more substances, without ceasing to exist, may be changed into one or more new and entirely different substances. When such an alteration occurs, it is called a chemical change, or interaction or reaction. To learn more about this most mysterious fact, we shall take up in the next chapter some examples of a simple and long familiar nature. Definition of Law. In science, a law, or generalization or rule, is a statement describing some general fact or constant mode of behavior. Its uses are to condense a great many similar facts into one statement, and thus to make the whole set of facts more easily memorable. Constituents. As we have seen, we speak of the substances in a mixture as the components. When we wish to refer to the forms of matter which are chemically united in a compound, we call them the constituents (Lat., standing together) of the com- pound substance. Thus, iron and oxygen are the constituents of rust. The chemist separates (p. 13) the components of a mixture, for that is all that is necessary. He liberates the constituents of a compound, however, because they are bound together in chemical combination. The names given to compounds are usually devised so as to indicate the nature of the constituents. Thus, iron-rust is oxide of iron (or ferric oxide, from Lat. ferrum, iron). The yellowish powder from lead is lead oxide or oxide of lead, and the white powder from tin is oxide of tin. SUBSTANCES AND PROPERTIES 9 A Condensed Form of Statement. We may represent a chemical combination, or indeed any kind of chemical change, in a condensed form, thus : Iron + Oxygen > Oxide of iron (ferric oxide). Each name stands for a substance. Two substances in contact with one another (mixed), but not united chemically, are con- nected by the + sign. The arrow shows where the chemical change comes in, and the direction of the change. We read the statement thus : Iron and oxygen brought together under suitable conditions undergo chemical change into oxide of iron, called also ferric oxide. Similarly we may write: Lead + Oxygen Oxide of lead. Tin + Oxygen > Oxide of tin. The Increase in Weight in Rusting. As we have seen, the process of rusting is accompanied by a slow increase in the weight of the solid, due to the gradual addition of oxygen to the metal. Now, this increase in weight ceases of its own accord, when a certain maximum has been reached. This occurs when the last particles of the metal have disappeared. Thus, the lead gains in weight until every 100 parts of the metal have gained 7.72 parts of oxygen, and the tin until every 100 parts have gained 26.9 parts of oxygen. When these increases have occurred, the metal is found to have been all used up, and prolonged heating and stirring cause no further union with oxygen and no further change in weight. This fact, that each substance limits itself of its own accord to combining with a fixed proportion of the other substance, in forming a given compound, is one of the most strik- ing facts about chemical combination. In mixtures, any propor- tions chosen by the experimenter may be used. In chemical union, the experimenter has no choice; the proportions are de- termined by the substances themselves. Thus, 100 parts of iron when turning into ordinary red rust take up 43 parts of oxygen, no more and no less. 10 TEXT BOOK OF CHEMISTRY This fact enables us to make our condensed statements more specific and complete by including in them the proportions by weight used in the chemical change: Iron (100) + Oxygen (43) -> Ferric oxide (143). Lead (100) + Oxygen (7.72) -* Oxide of lead (107.72). The following numbers, which represent the same proportions by weight, are the ones commonly used by chemists: Iron (111.68) + Oxygen (48) -> Ferric oxide (159.68). Summary. Thus far, we have learned that chemistry deals with substances and their physical properties, and with the changes which substances undergo. We have discussed and defined a number of important words expressing fundamental chemical ideas. Finally, we have touched upon the weights of the ma- terials used in chemical change, a subject of great importance which will be more fully developed in a later chapter. Exercises. 1. Describe the following, by mentioning some of their specific properties: (a) water, (b) wool, (c) cotton (pp. 2-3). 2. If any of the following are mixtures, mention the facts which show them to contain more than one substance (p. 13) : (a) muddy water, (b) an egg, (c) milk. 3. State and illustrate the first two laws of chemistry (pp. 5-7). 4. Make definitions of "pure" and " impure" as applied to a sample of a substance (p. 6). 5. Give a list of the specific properties mentioned in this chapter. 6. In recognizing a specimen to be quartz, does the chemist consider (a) the weight, (b) the temperature, (c) the length of the specimen (p. 4) ? If not, why not? 7. Take one by one the words or phrases printed in black type and the titles of the sections in this chapter, and endeavor to recollect what you have read about each. In each case try, (a) to recall the meaning and to state it in your own words; (b) to recall SUBSTANCES AND PROPERTIES 11 the facts associated with, and the reasoning which lead up to the point in question; (c) to recall examples illustrating the concep- tion and to apply the conception in detail to each example. When- ever memory fails to give a perfectly clear report of the matter in hand, the text must be read and re-read until the essential point can be repeated from memory. Use the same method in all future chapters. A useful prac- tice is to employ a pencil as you read and to underline systemati- cally all the important facts and statements, and then to go back and apply to each marked place the process described above. 8. Define the following terms : Specific gravity, tenacity, melt- ing-point, specific physical property, pure body, vacuum. 9. Is it logical to say "pure substance?" 10. Why do we decide that granite is a mixture and iron a cingle substance? 11. What weight of oxygen would be required to convert 25 grams of lead into oxide of lead? 12. Make a list of the technical words we have defined, and place the definition opposite to each. CHAPTER II CHEMICAL CHANGE AND THE METHODS OF STUDYING IT WE must now take up two new examples of chemical change. They will aid us in introducing one or two additional conceptions and laws. These are continually used by the chemist, and without them we cannot begin the systematic study of the science. Another Case of Combination: Iron and Sulphur. Since oxygen is an invisible gas, there is a slight difficulty in real- izing that rusting consists in the union of two substances this gas and a metal. The present example is less interesting histori- cally, but it is simpler because both substances are visible and are easily handled. The case of iron and sulphur will enable us to illustrate the same point of view and to practice the application of the same technical words. It will also introduce us to two manipulations filtration and evapora- tion which are frequently used by the chemist. We begin by observing the physical prop- erties of the two substances. Sulphur is a pale-yellow substance of low specific gravity (sp. gr. 2). It is easily melted (m.-p. 112.8 C.). It does not dissolve in water that is, it does not mix completely with and disappear in water, as sugar does on stirring. It does dissolve readily in carbon disulphide, however. It crystallizes in rhombic forms (Fig. 7). It is not attracted by a magnet. Study of the Mixture, before Combination. Now, if gome iron filings and pulverized sulphur are stirred together in a 12 CHEMICAL CHANGE AND METHODS OF STUDYING IT 13 FIG. 8. mortar, the result is a mixture. True, the color is not that of either substance, but with a lens particles of both substances can be seen. Passing a magnet over the mixture will easily remove a part of the iron, and with the help of a lens and a needle the mixture could be picked apart particle by particle, com- pletely. We can separate the components of the mixture more expeditiously, however, by using manipulations based upon certain suitable properties. Thus, sulphur dissolves in carbon disulphide while iron does not. If, therefore, a part of the mixture is placed in a dry test-tube along with some carbon disulphide (Fig. 8), and is shaken, the liquid dissolves the sulphur and leaves the iron. To complete the separation, the iron must be removed from the liquid by filtration, and the sulphur recovered by evaporation of the carbon disulphide. Filtration. Iron, or any solid, when it is mixed with a liquid or with a solution (like the solution of sulphur in carbon disulphide) is said to be sus- pended in the liquid. If the solid is one that settles rapidly, the liquid may be separated from the solid, in a rough way, by pouring off as much of the clear, supernatant liquid as possible. This is called decantation. A complete separation is effected by pouring the mixture on to a cone of filter paper supported in a glass funnel (Fig. 9). The liquid, together with anything that may be dissolved in it, runs through the pores of the paper and down the hollow stem of the funnel. The liquid is then called the filtrate. The particles of the suspended solid are too large to pass through the pores, and so collect on the surface of the filter FIG. 9. 14 TEXT BOOK OF CHEMISTRY paper. This operation, like everything the chemist does, takes advantage of the physical properties of the various materials. The material remaining on the paper (the residue), when dry, is wholly attracted by a magnet and shows all the other properties of iron. Evaporation. To recover the sulphur, the solution in carbon disulphide the filtrate is poured into a porcelain evaporat- ing dish (Inflammable! Keep flames away). When the vessel is set aside, the liquid gradually passes off in vapor (e-vapor-ates) . Sulphur, however, does not evaporate at room temperature and remains as a residue, in the form of crystals of rhom- bic outline in the bottom of the dish (Fig. 10). Here, again, physical proper- ties have been utilized. Since the physical properties of two substances are not changed by mixing, we have thus used the properties of the iron and sulphur so as to separate them once more. The iron is on the paper; the sulphur is in the dish. Combination of Iron and Sulphur. But iron and sulphur are capable of combining. If we alter the conditions by raising the temperature of some of the dry mixture, as we did in causing lead to rust rapidly, chemical union sets in. When we place some of the original mixture of iron and sulphur into a clean test-tube and warm it, we soon notice a rather violent development of heat taking place, the contents begin to glow, and what appears to be a form of combustion spreads through the mass. The heating em- ployed at the start falls far short of accounting for the much greater heat produced. When these phenomena have ceased, and the test-tube has been allowed to cool, we find that it now contains a somewhat porous-looking, black solid. This material is brittle; it is not magnetic; it does not dissolve in carbon di- CHEMICAL CHANGE AND METHODS OF STUDYING IT 15 sulphide; and close examination, even under a microscope, does not reveal the presence of different kinds of matter. This sub- stance is known to chemists as ferrous sulphide and, as we see, its properties are entirely different from those of its constituents. In this connection we must not omit to notice that, as in rusting, a certain fixed proportion will be used in forming the compound. We find that, for 7 parts of iron, almost exactly 4 parts by weight of sulphur are required. If more iron is put into the original mixture, then some unused iron will be found in the mass after the action. If too much sulphur is employed, some may be driven off as vapor by the heat and any that remains, beyond the correct proportion, can be dissolved out of the ferrous sulphide with car- bon disulphide. The sulphur which has combined with the iron, however, is no longer present as sulphur it has no longer the properties of sulphur, and therefore cannot be dissolved out: Iron (55.84) + Sulphur (32.07) - Ferrous sulphide (87.91). Another Illustration: Mercuric Oxide. It has long been known that air contains an active and an inactive gas. The Chinese called them yin and yang, respectively. Mayow (1643- 1679) showed that the active gas caused rusting, that it was absorbed by paint (really by the linseed oil) in "drying," that it supported com- bustion of wood and sulphur, and that it is nec- essary to life, being absorbed by the blood from the air entering the lungs. It-was not until 1774, however, that a pure specimen of this gas was obtained, by Bayen, and was recognized to be FI 11 a special kind of gas different from ordinary air. The gas (later to be named oxygen) was made by Bayen from mercuric oxide, a bright red, rather heavy powder. When the oxide is heated (Fig. 11), we find that a gas is given off. This gas is easily shown to be different from air, since a glowing splinter of wood is instantly relighted on being immersed in it. The gas is 16 TEXT BOOK OF CHEMISTRY pure oxygen. During the heating, we notice also that a metallic coating appears on the sides of the tube, in the form of a sort of mirror. Apparently the vapor of some metal is coming off with the oxygen and condensing on the cool parts of the tube. As this shining substance accumulates it takes the form of globules, which may be scraped together. It is, in fact, the metal mercury, or quicksilver. If the heating continues long enough, the whole of the red powder eventually disappears, and is converted into these two products. Second Variety of Chemical Change: Decomposition. Bayen's experiment introduces to us a second, and very common kind of chemical action. The first variety was combination or union (p. 14). The second is called decomposition. It consists in starting with a single substance (here mercuric oxide) and splitting it into two (or more) substances, which differ in properties from the substance taken and from one another. Here, the red powder gave mercury, a liquid metal, and oxygen, a colorless gas. Simple and Compound Substances. We have seen that two (or more) substances, like lead and oxygen, can combine to form a compound substance. Are all substances, then, com- pounds? We find that some are not. We have never succeeded in obtaining lead, or oxygen, or iron, or tin, or sulphur by com- bining any two substances. We can decompose mercuric oxide by heat, and we have other ways of decomposing compounds like oxide of tin and ferrous sulphide, but we have never succeeded in decomposing the mercury or the oxygen, the iron or the sulphur themselves. Substances which we are not able, at will, to decom- pose into, or to make by chemical union from, other substances are called simple or elementary substances. The distinction be- tween simple and compound substances was first drawn by Lom- onossov in 1741. Later, and independently, it was stated very clearly by Lavoisier (1789). CHEMICAL CHANGE AND METHODS OF STUDYING IT 17 Several substances, regarded in Lavoisier's time as elementary, have since been shown to be compounds. Thus, quicklime was a simple substance until Davy, in 1808, prepared the metal calcium and showed that quicklime was the oxide of this metal. Hence, we do not say that the substances regarded as simple cannot be decomposed, but only that they are substances which we "are not able" (at present) to decompose. The phrase "at will" is also important. Radium (q.v*) cannot be decomposed at will, but it undergoes continuous "disintegra- tion," producing the elements helium and lead. We can neither hasten, retard, nor stop this spontaneous decomposition. Elements. The word element is used in two senses. It is applied to the simple substance. Thus we speak of "the element iron," meaning the metal iron. It is applied also to the iron- matter contained in ferrous sulphide or in ferric oxide. The reader should note that it is correct usage to speak of the element iron and the element sulphur in ferrous sulphide, but a chemist would never say that this compound contained the simple sub- stances iron and sulphur. If he did, we should understand him to mean that it was a mixture, and we should expect parts of the material to be magnetic like iron, and other parts to be yellow and soluble in carbon disulphide, which is not the case. In the same way the name of an element (such as iron) is applied both to the material in combination and to the free substance. Thus "iron" may mean free, uncombined, metallic iron, or iron-matter in some compound. The sense in which the word is employed must be inferred from the context or circumstances. When a chemist speaks, as he sometimes does, colloquially, of "iron" in a drinking water, for example, we know at once that he refers to iron in the form of some compound, for metallic iron does not * Contraction for quod vide, which see. This abbreviation is used when subjects not yet discussed are mentioned. For such subjects, consult the index. 18 TEXT BOOK OF CHEMISTRY dissolve in water and, if it did, would quickly turn into rust or some other form of combination. The word element, then, means one of the simple forms of matter, either free or in combination. In formally describing a body or specimen, the chemist always avoids the ambiguity just referred to by naming the components, i.e., the substance or substances it contains. He assumes that the nature and constituents of these substances will be known to anyone hearing or reading the description. If he says the body con- tains zinc and sulphur, it is understood that the body is a mixture of these simple substances. If it contained these elements in combination, the chemist would report that it was sulphide of zinc. cc. The Common Elements. Thousands of dif- ferent compound substances are known but, when they are decomposed, it is found that the number of different elements contained in them is not great. Dozens of substances contain iron, hundreds con- tain sulphur, thousands contain oxygen. In fact, by combining a limited number, two, three, or four, of simple substances together, in varying propor- tions by weight, an almost unlimited number of different compound substances could be produced. A list of the elements appears on the inside of the cover, at the end of this book, and contains about eighty names. Of these, a large number are rare, and seldom encountered. More than 99 per cent of terrestrial material is made up of eighteen or twenty elements and their compounds. Only about twenty elements occur in nature in their simple, uncombined condition. Three-fourths of the whole number are found in combination exclusively, and must be liberated by some chemical action. HYO. TIT. "CHLOK. OTHCH ELTS. FIG. 12. CHEMICAL CHANGE AND METHODS OF STUDYING IT 19 Taking the atmosphere, all terrestrial waters, and the earth's crust, so far as it has been examined, F. W. Clarke has estimated the plentifulness of the various elements. The first twelve, with the quantity of each contained in one hundred parts of terrestrial matter, and constituting together 99 per cent, are as follows: Oxygen . . Silicon . . . 49.85 . 26.03 Calcium . . . Sodium . . , . 3.18 . 2.33 Hydrogen . Titanium . . 0.97 . 0.41 Aluminium . 7.28 Potassium . . , . 2.33 Chlorine . . . 0.20 Iron . . . . 4.12 Magnesium . . . 2.11 Carbon . . . 0.19 Thus oxygen accounts for nearly one-half of the whole mass. Silicon, the oxide of which when pure is quartz and in less pure form constitutes ordinary sand, makes up half of the remainder. Valuable and useful elements, like gold, silver, sulphur, and mer- cury, are among the less plentiful which, all taken together, furnish the remaining one per cent. Law of Definite Proportions. In the decomposition of mercuric oxide (p. 15) * we find that, for every 100 parts of mer- cury liberated, almost 8 parts of oxygen (more exactly 7.97 parts) by weight are set free. Using the numbers commonly employed in chemistry, which represent the same proportion by weight : Mercuric oxide (216.6) -> Mercury (200.6) + Oxygen (16). We find also that mercury and oxygen can be made to combine to form mercuric oxide, and the proportions by weight required are the same. Moreover, every sample of mercuric oxide, whether made by combination, or in any of the other possible ways, always contains this proportion of the two elements. We have already * References to previous pages are used in order to save needless repetition in writing. The beginner requires endless repetition in his reading, however, and must form the habit of examining, in conjunction with the current text, the parts referred to. The passages cited are, by the reference, made part of the current text, which will usually not be clear without them. The same remark applies to topics referred to by name. Such topics must be sought in the index. All terms, and especially those borrowed from physics, if not perfectly familiar, must be looked up in a work on physics or in a dictionary. 20 TEXT BOOK OF CHEMISTRY seen that the oxides of lead and tin contain fixed proportions (p. 10) of the metal and oxygen and that ferrous sulphide has a constant composition by weight. The same principle is found to apply to all chemical compounds, and is stated in the law of definite or constant proportions : In every sample of any compound substance, formed or decomposed, the proportion by weight of the constituent elements is always the same. (For the only known exception to this law, see radium.) Conservation of Mass. The most painstaking chemical work seems to show that, if all the substances concerned in a chemical change are weighed before and after the change, there is no evidence of any alteration in the quantity of matter. The two weights, representing the sums of the constituents and of the prod- ucts, respectively, are, indeed, never absolutely identical, but the more careful the work and the more delicate the instrument used in weighing, the more nearly do the values approach identity. We are able to state, therefore, that the mass of a system is not affected by any chemical change within the system. This statement simply means that the great law of the conserva- tion of mass holds true in chemistry as it does in physics. Chemi- cal changes, thoroughgoing as they are in respect to all other qualities, do not affect the mass; an element carries with it its weight, entirely unchanged, through the most complicated chemi- cal transformations. Superficial observation, as of a growing tree, might seem to give evidence of the very opposite of conservation of matter. But here the carbon dioxide gas in the air, the most important source of nourishment for plants, is overlooked. Similarly, the gradual disappearance of a candle by combustion seems to illustrate the destruction of matter. But if we catch the gases which rise through the flame (Fig. 13), we find that the gases weigh even more than the part of the candle which has been sacrificed in making them. When we take account of the weight of the oxygen obtained from CHEMICAL CHANGE AND METHODS OF STUDYING IT 21 the air which sustains the combustion, we find that there is really neither loss nor gain in weight. If we carry out chemical changes in closed vessels (Fig. 14), which permit neither escape nor access of material, we find that the weight does not alter. FIG. 13. FIG. 14. Physics in Chemistry. It will be seen that one cannot accomplish anything in chemistry without acquiring and using some knowledge of physics. We measure quantities by means of the physical attributes, weight and volume. We produce chemi- cal change by arranging the physical conditions, for example, by mixing, heating, or using an electric current. Physical means are the only means we possess for producing, stopping, or modifying chemical changes. Again, we ascertain whether a chemical change has taken place or not by observing the physical properties of the materials before and after the experiment. Thus, we noted that the red, powdery oxide of mercury, when heated, gave a liquid metal and a gas. All the phenomena of chemistry are physical. A phenomenon is literally something that is seen or, more generally, something that affects any of the senses. Observ- ing physical phenomena is, therefore, our sole means of studying chemical changes. Chemical work is, in fact, entirely dependent 22 TEXT BOOK OF CHEMISTRY upon the skilful use of physical agencies, and upon the close obser- vation of physical phenomena for its success. It is only the inference, following the experiment and the obser- vation, that is strictly chemical. If one substance gives two different substances, or if two substances give one different sub- stance, for example, we infer that a chemical change has occurred. We then try to recognize the substances by their properties and name them. Changes like that of ice into water, or of water into steam, and vice versa, are not regarded as chemical changes. These are called changes of state, or of state of aggregation. The solid, liquid, and gaseous forms are different states of the same substance. Law: Explanation: Scientific Method. There is a widely spread impression that a science, like chemistry, is a part of the natural order of the universe. It is thought that we are try- ing to find the boundaries of chemistry, as they have been pre- determined by nature, and to discover the facts, relations of facts, and laws which nature has provided as a means of classifying the content of the science. Now, the situation is precisely the re- verse of this. Nature provides only the materials and the phe- nomena, and man is attempting to classify them. He divides the whole into groups, such as physics, chemistry, botany, etc. Then he classifies the facts within each group, in order that he may more easily remember them and perceive their relations. He often finds that, when new facts are discovered, parts of the classification have to be changed. In the preceding pages, we have discussed some of the ways that have been invented for classifying the materials and facts assigned to chemistry. Thus, we pick out a number of facts of a like nature and try to make a single statement which will cover all these facts. For example, we find about one hundred thousand different substances and, in the case of each substance, every speci- men that we have examined contains the same proportions of CHEMICAL CHANGE AND METHODS OF STUDYING IT 23 the constituent elements. So we formulate the law of constant proportions. A law or generalization in chemistry is a brief state- ment describing some general fact or constant mode of behavior. We must remember, however, that laws are only true so long as no facts in conflict with them are known. There are no laws in nature. Nature presents materials and phenomena as she pleases. The laws are parts of science, which is made by man, and is a description of natural facts as man knows them. As we have seen (p. 17), at least one undoubted exception to the law of constant propor- tions has recently (1914) been discovered, and other exceptions to this law will undoubtedly present themselves. One section (p. 9) was entitled: " Explanation of rusting." If that paragraph be now re-read, it will be found that, in the ordinary (as distinct from the scientific) sense of the word, no explanation was given! When we ask a man to " explain" some feature in his conduct, we recognize that he might have chosen to act otherwise, and we wish to know why he acted precisely as he did. Nature, however, has no free will, and cannot tell why she presents certain phenomena, and not others. On examining the explanation, we find that it simply shows that when iron rusts it combines with oxygen from the air. This is an additional fact. It shows how iron rusts, namely, by taking up oxygen, but not why it is able to unite with oxygen. We simply do not know why iron can combine with oxygen gas and platinum cannot. Explanations in chemistry are of three kinds. (1) We usually try to show that the phenomenon is not an isolated one. Thus, we show that other metals rust. This reconciles us to some ex- tent to the fact that iron rusts, and we feel some mental satisfac- tion. This is the method of showing that the fact to be explained is a member of a large class of similar facts. (2) Next, we try to get more information about the fact to be explained. Thus, when, to the acquaintance with the outward manifestations of rusting, we add the further information that there is an increase in weight, 24 TEXT BOOK OF CHEMISTRY and that this is due to union of oxygen from the air with the iron, we feel increased satisfaction, and say that the fact has been "ex- plained." (3) If we are still dissatisfied, and can discover no further useful facts, we imagine a state of affairs which, if true, would classify the fact or add to what we know about it. This step we call explaining by means of an hypothesis. We then devote our attention to trying to verify the hypothesis. The formulation of laws and the making of attempts to explain facts are part of what is called the scientific method. The purpose of this method is to convert the subject matter into a science, that is, into an organized body of knowledge. Summary. In this chapter we have learned: (1) that, while there are many substances, there is a limited number of entirely different kinds of matter (elements) ; (2) that, in addition to constant physical properties, each substance has a constant composition by weight. We have also learned that physical properties are utilized in manipulations, like filtration and evap- oration, as well as for identifying substances, and that physical attributes are used for measuring quantities in chemistry and physical conditions for guiding chemical change. Finally, we have seen that a science is not a natural, but a manufactured product, and that the science of chemistry is still in the making. Exercises.* 1. What physical properties are used (a) in filtration, (b) in evaporation, (c) in the separation and identifica- tion of the products from heating mercuric oxide (p. 15) ? 2. Describe: (a) a red-hot rod of iron, 10 cm. long by 1 cm. diameter, weighing 58.5 g.; (6) a solution of 5 g. of sulphur in 20 c.c. (59 g.) of carbon disulphide at 18 C. In doing so, divide the description into attributes, conditions, and properties. * The exercises should in all cases be studied with minute care. They not only serve as tests to show that the chapter has been understood, but very frequently (as in No. 4) also call attention to ideas which might not be ac- quired from the text alone, or (as in Nos. 1, 2, 5) assist in elucidating ideas given in the text which, without the exercises, might not be fully grasped. CHEMICAL CHANGE AND METHODS OF STUDYING IT 25 3. Consider the following materials and state whether, so far as you can now judge, each is a single substance or a mixture: (a) a candle, (b) a cake of soap, (c) an egg. 4. What are the two most direct ways of showing a substance to be a compound ? Illustrate each. 5. If we say that quicklime contains calcium, do we mean the element or the simple substance calcium? 6. What explanation was given, (a) of the disappearance of mercuric oxide when heated, (b) of the absence of iron and sulphur, as substances, from ferrous sulphide? Which of the three kinds of explanation was used in each case ? 7. What weight of oxgyen will be required to combine with 15 grams of lead (p. 10)? 8. If 5 grams of lead and 4 grams of oxygen were heated together, which of the two would remain in part unused? How much of this one would remain (p. 10) ? CHAPTER III AIR AND OXYGEN WE have seen that metals absorb a gas, called oxygen, from the air, and turn into a rust or oxide. Let us now consider what happens to the air during this process. The Nature of Air. Can a metal, like iron, when rusting, absorb the whole of a sample of air, or does it select a part of the air only? If we sprinkle some powdered iron in a test-tube, having first moistened the interior to cause the powder to adhere to the inside surface, and then set the tube, mouth downwards, in a dish of water (Fig. 15), we obtain be- fore long an answer to this question. As the iron slowly removes the oxygen, the pressure of the atmosphere outside pushes the water up the tube. But, after ascending only about one-fifth of the total height of the tube, the water comes to rest. Inspection shows reddened particles where rusting has taken place, but much of the iron is still dark grey, and is as little able to rust in the remaining gas as in a vacuum. Four-fifths of the air, then, is composed of gases which do not combine with iron, and only one-fifth is oxygen. The four- fifths is in fact almost all (99 per cent) nitrogen, a substance which combines with very few materials, while the balance (1 per cent) is made up of gases which do not enter into combination with any known substances. Oxygen, on the contrary, combines with almost all simple substances, although in many cases, as in those of lead and tin, heating is required to hasten the process. 26 FIG. 15. AIR AND OXYGEN 27 Activity and Stability. A substance which enters into combination easily, is called active, so that oxygen is spoken of as an active element, nitrogen as a relatively inactive or indifferent element. An active element, since it combines greedily, naturally holds tenaciously to that with which it has combined. An active element means, therefore, also, one which is in general difficult to liberate from combination. Its compounds are in general rela- tively stable. Law of the Influence of Heating. Even oxygen, active as it is, does not, visibly, combine with tin, when both are cold. Lead rusts very slowly at the ordinary temperature. Iron rusts very much faster when heated than when cold. In every chemical change we find that raising the temperature hastens the process. Other things being equal, it causes a greater quantity of material to undergo the change in a given time. This is the third law of chemistry. Oxygen. We cannot do better than begin the more systematic study of chemistry with oxygen, for it is a most interesting as well as useful substance. It is the active component of the air. We depend upon it for life, since in its absence we suffocate, for heat, since wood, coal, and gas will not burn without it, and even for light where oil, gas, or a candle is used. We wish to know with which substances we use in the laboratory it can combine, as well as the substances on which it has no action. This information will show us how to work, m future, with- out interference from the oxygen in the air and whether oxygen has probably played a part in some experiment or not. Let us take up, then, (1) the history of the element, (2) what materials contain oxygen (occurrence), (3) how we can obtain it in a pure state (preparation), (4) what its specific physical properties as a substance are, and (5) what it does, and what it cannot do in nature and in the laboratory (chemical properties). The classifi* 28 TEXT BOOK OF CHEMISTRY cation of the facts about this, and other substances under five heads, is somewhat mechanical, but has the advantage of enabling the reader quickly to find any required information. History of Oxygen. The Chinese, in or before the eighth century, knew that there were two components in the air, and that the active one, yin, combined with some metals, and with burning sulphur, and charcoal. They even knew that it could be obtained in pure form by heating certain minerals, of which one was saltpeter. Leonardo da Vinci (1452-1519) seems to be the first European to mention the presence of two gases in the air. Mayow (1669) measured the proportion of oxygen in the air and discussed fully its uses in combustion, rusting, vinegar-making, and respiration, but did not make a pure sample. Hales (1731) made it from saltpeter, and measured the amount obtainable, but did not see any connection between it and the air! Bayen (Apr., 1774) was the first to make it by heating mercuric oxide. Priestley (Aug. 1, 1774) made it by heat- ing the same substance and quite pur- poselessly, as he admits, thrust a lighted candle into it and was delighted with the extreme brilliance of the flame. He had, however, entirely incorrect ideas about FlG 16 its nature, and described it as unbreak- able and noxious (poisonous). Scheele, a Swedish apothecary, had made it in 1771-2 from no less than seven different substances and understood clearly that atmos- pheric oxygen combined with metals, phosphorus, hydrogen, linseed oil and many other substances. But the publisher did not get his book out until 1777, and Priestley is usually, though incorrectly, credited with the " disco very" of the element. Finally, Lavoisier (1777) heated mercury in a retort (Fig. 16), the neck of which projected into a jar standing in a larger dish of mercury. The air, thus enclosed within the jar and English physician and chemist, 1643-1679 AIR AND OXYGEN 29 the retort, during twelve days lost one-fifth of its volume. Sim- ultaneously, red particles of mercuric oxide accumulated on the surface of the mercury in the retort. The residual gas, nitrogen, no longer supported life or combustion. The oxide, on being heated more strongly, by itself, gave off a gas whose volume ex- actly corresponded with the shrinkage undergone by the enclosed air, and this gas possessed in an exaggerated degree the properties which the air had lost. The proof that oxygen was a component of the atmosphere was therefore complete. Later, Lavoisier, in the mistaken belief that the new element was an essential con- stituent of all sour substances, named it oxygen (Gk, add-producer) . Occurrence. As we have seen, nearly 50 per cent of terres- trial matter is oxygen. Water contains about 89 per cent, the human body over 60 per cent, and common materials like sandstone, limestone, brick, and mortar more than 50 per cent of this element. One-fifth by volume (nearly one-fourth by weight) of the air is free oxygen. Preparation of Oxygen. 1. The oxygen of com- merce is now made chiefly from liquefied air (q.v*). The liquid oxygen boils at 182.4, but the nitrogen boils at an even lower temperature (194). Since the liquid air has a temperature of about 190, somewhat above that of boiling nitrogen, the latter evaporates much more freely than does the oxygen. After a time, when the remaining liquid is almost pure oxygen (96 per cent), the gas coming off is compressed by pumps into the steel cylinders (Fig. 17) in which it is sold. In medicine, patients suffering from pneumonia or suffocation obtain some relief by inhaling it in this form. It is also used in feeding flames, instead of air, when intense heat is required (see acetylene torch and calcium light). See p. 17, footnote 30 TEXT BOOK OF CHEMISTRY 2. Unfortunately, it is difficult to liberate oxygen from natural substances. Saltpeter (potassium nitrate), for example, which is found in many soils and can be dissolved out with water, gives off oxygen only when raised to a bright red heat by the Bunsen flame or blast lamp. But, even at this temperature, it gives up only one-third of the oxygen it contains. 3. In practice, we are compelled to use manufactured sub- stances. Amongst the artificial substances are mercuric oxide, expensive but historically interesting (p. 15), potassium chlorate, perhaps the most convenient for laboratory use, and sodium per- oxide. Potassium Chlorate (q.v.) is a white crystalline substance used, on account of the oxygen it contains, in large quantities in the manufacture of matches and fireworks. When heated in a tube similar to that in Fig. 11, it first melts (334) and then, on being more strongly heated, it effervesces and gives off a very large volume of oxygen. Examination shows that the whole of the oxygen it contains (39 per cent) can be driven out. The white FIG. 18. material which remains after the heating is identical with the mineral sylvite. To the chemist it is known as potassium chloride. The change, together with the weights of the materials, is as follows: Potassium chlorate (122.56) > Potassium chloride (74.56) + Oxygen (48) Potassium (39.1) Chlorine (35.46) Oxygen (48) Potassium (39.1) Chlorine (35.46) AIR AND OXYGEN 31 A peculiarity of this action is that admixture of manganese dioxide (the mineral pyrolusite) increases very markedly the speed with which the decomposition of the potassium chlorate takes place. Hence, in its presence, and it is generally mixed with the chlorate in laboratory experiments (Fig. 18), a sufficient stream of the gas is obtained at a relatively low temperature (below 200, see p. 28). Hales (p. 28) was the first to collect a gas over water (Fig. 18), in order that it might be kept unmixed with air. 4. Oxygen can be obtained conveniently from sodium peroxide and water by means of generators (Fig. 19) similar to the acety- lene generators used on automobiles. When the metal sodium is burned in air, sodium peroxide is obtained as a powder. This powder, after being melted, solidifies in com- pact, solid form, and is sold as oxone. The oxone is bought in a small, sealed tin can, the ends of which are perforated in several places just before use. When the valve (B) is opened, so that the oxygen escapes, the water, which fills the generator almost to the top, enters the can (C) by the holes in the bottom and interacts with the oxone. When the valve is shut, the gas continues to be generated until it has driven the water down again below the level of the bottom of the can. Sodium peroxide (78) + Water (18) -> Sodium hydroxide (80) + Oxygen (16) Sodium (46) Hydrogen (2.016) Sodium (46) Oxygen (32) Oxygen (16) Oxygen (32) Hydrogen (2.016) This method is convenient because it works at room temperature and can be started and stopped at will. The sodium hydroxide produced is very soluble in water and remains dissolved. Note that the name of this substance indicates the elements which compose it. 7/ FIG. 19. 32 TEXT BOOK OF CHEMISTRY Catalytic or Contact Action. The influence of manganese dioxide in causing the potassium chlorate to decompose more easily (p. 30) well deserves notice. The effect is very striking if some pure potassium chlorate is melted carefully, to avoid superheat- ing, in a wide-mouth flask (Fig. 20). The flask is provided with a wide exit tube, from which a rubber tube may lead to a bottle inverted in a trough filled with water as in Fig. 18. A little manganese dioxide is contained in the upper, closed tube. No effervescence of the chlorate can be seen at its melting- point (334) only a little air, expanded by the heating, issues from the tube. When, how- ever, the closed tube contain- ing the manganese dioxide is rotated into a vertical position (see dotted lines), and the black powder falls into the chlorate, the oxygen comes off in tor- rents, in consequence of the enormous acceleration of the decomposition. As a precaution against injury from an explosion of the flask, it is advisable to wrap the latter in a towel before turning the tube. It must also be noted that the manganese dioxide is not itself permanently altered. If the material left after the action is shaken with water, the potassium chloride dissolves, while the dioxide does not. Filtration (p. 13) thus enables us to recover the latter, and to ascertain that it has been changed neither in quantity nor in properties. The only effect of the dioxide is to hasten the decomposition of the chlorate, which would otherwise be too slow at 200 (p. 30), or even at 334 (its m.-p.) to be of any practical value. Sub- stances which hasten a chemical action without themselves under- FIG. 20. AIR AND OXYGEN 33 going any permanent change are called contact agents, catalytic agents, or catalysts. The process is called contact action or catal- ysis (Gk., decomposition, not a very fortunate choice of words). Such substances are frequently used in chemistry. The addition of a suitable catalyst is one of the conditions (p. 32) for carrying out actions in which a contact agent is necessary. Many sub- stances of this class are secreted by animals and plants and play an important part in digestion, fermentation, and other physiologi- cal changes. Their presence often enables very complex chemi- cal actions to proceed rapidly at rather low temperatures. The oxone, mentioned above, always contains a trace of cuprous oxide 4Cu20,H 2 O which hastens the action on water. Specific Properties of Two Kinds, Physical and Chemical. We have learned that every substance has its own set of specific properties. In describing a substance, it is convenient to divide the properties into two classes. The list of substances with which the given substance can enter into chemical combination, for example, we place under specific chemical properties. Relations of the substance to any of the varieties of chemical change belong to this class. On the other hand, we do not consider melting or boiling to be chemical changes, so we place the temperatures at which the sub- stance melts (m.-p.) and boils (b.-p.), its color, etc. (for list, see p. 7), under specific physical properties. Properties of either class may be used for recognizing a substance. Specific Physical Properties of Oxygen. Oxygen resembles air hi having neither color, taste, nor odor. The density of a sub- stance is, strictly speaking, the weight of 1 cubic centimeter (1 c.c.). In the case of a gas, we frequently prefer to give the weight of 1000 c.c. (1 liter), at and 760 mm. (1 atmosphere) barometric pressure. For oxygen this weight is 1.42900 grams (Morley). The corre- sponding weight for air is 1.293, so that oxygen is slightly heavier, 34 TEXT BOOK OF CHEMISTRY bulk for bulk, than air (in the ratio 1.105 : 1). Oxygen can be liquefied by compression, provided its temperature is first reduced below 118, which is its critical temperature.* The gas is slightly soluble in water, the solubility at being 4 volumes of gas in 100 volumes of water (at 20, 3 : 100). The solubility of oxygen in water, although slight, is in some respects its most important physical property. Fish obtain oxy- gen for their blood from that dissolved in the water. With air- breathing animals (like man), the oxygen could not be taken into the system, if it did not first dissolve in the moisture contained in the walls of the air sacs of the lungs, and then pass inwards in a dissolved state to the blood. Liquid oxygen, first prepared by Wroblevski, has a pale-blue color. At one atmosphere pressure, that is, in an open vessel, it boils at 182.5. Its density (weight of 1 c.c.) is 1.13, so that it is slightly denser than water. By cooling with a jet of liquid hydrogen, Dewar froze the liquid to a snow-like, pale-blue solid. A tube of liquid oxygen is noticeably attracted by a magnet. Six Specific Physical Properties of Each Gas. Although every substance has many physical properties, we shall mention only those which are used in chemical work, with occasionally the addition of any peculiar or unexpected quality. It will aid the memory to recall the physical properties of a gas, if we note that, as a rule, only six such properties are mentioned: (1) color, (2) taste, (3) odor, (4) density, (5) liquefiability, defined by the critical temperature, (6) solubility, usually in water only. Specific Chemical Properties of Oxygen. The chemical properties of pure oxygen are like those of atmospheric air, only more pronounced. * Each gas has an individual critical temperature (q.v.) above which no pressure, however great, will produce liquefaction. The farther the tempera- ture of a specimen of the gas is below the critical point, the less will be the pressure required to liquefy it. AIR AND OXYGEN 35 Non-metallic Elements. Sulphur, when raised in advance to the temperature necessary to start the action, unites vigorously with oxygen, giving out much heat and producing a familiar gas hav- ing a pungent odor (sulphur dioxide). This odor is frequently spoken of as the " smell of sulphur," but in reality sulphur itself has no odor, and neither has oxygen. The odor is a property of the compound of the two. The mode of experimentation can be changed and the oxygen led into sulphur vapor through a tube. The oxygen then appears to burn with a bright flame, giving the same product as before. Phosphorus, when set on fire, blazes in oxygen very vigorously, forming a white, powdery, solid oxide phosphorus pentoxide. Burning carbon, in the form of charcoal or hard coal, glows bril- liantly and is soon burnt up. It leaves an invisible, odorless gas carbon dioxide. At high temperatures, oxygen combines readily with one or two other non-metals (e.g., silicon, boron, and arsenic), and to a small extent (1 per cent at 1900) with nitrogen. It will not combine directly with chlorine, bromine, or iodine, although oxides of the first and last can be prepared by using other varieties of chemical change. With the six members of the helium family (q.v.), of which no compounds are known, and with flourine, oxygen forms no compounds. Sulphur (32.06) + Oxygen (32) -^Sulphur dioxide (64.06). Phosphorus (62.08) +Oxygen (80) ^Phosphorus pentoxide (142.08). Carbon (12)+ Oxygen (32) ^Carbon dioxide (44). Metallic Elements. Iron, as we have seen, rusts exceedingly slowly in air and, even when red-hot, gives hammer-scale, the black solid which is broken off on the anvil, rather deliberately. In pure oxygen, a bundle of picture-wire, if once ignited, will burn with surprising brilliancy, throwing off sparkling globules of the oxide, melted by the heat. This oxide is a black, brittle substance, identical with hammer-scale, and different from rust (ferric oxide). 36 TEXT BOOK OF CHEMISTRY It contains, in fact, a smaller proportion of oxygen than does the latter, and is called magnetic oxide of iron. Iron (167.52) + Oxygen (64) -> Magnetic oxide of iron (231.52). All the familiar metals, excepting gold, silver, and platinum, when heated, combine with oxygen, some more vigorously, others less vigorously than does iron. Oxides of the three metals just named can also be made, but only by varieties of chemical change other than direct combination. Compound substances, if they are composed largely or entirely of elements which combine with oxygen, are able themselves to interact with oxygen. Usually, they produce a mixture of the saiLe oxides which each element, separately, would give. Hence, wood, which is composed of carbon and hydrogen with some oxygen, when burnt in oxygen, produces carbon dioxide and water (oxide of hydrogen) in the form of vapor. Again, carbon disul- phide burns readily, giving carbon dioxide and sulphur dioxide, just as do carbon and sulphur, separately. Ferrous sulphide gives, similarly, sulphur dioxide and magnetic oxide of iron. Tests. A Test for Oxygen. A test is a property which, because it is easily recognized (a strong color, for example), or for some other sufficient reason, is commonly employed in recognizing a substance. Oxygen, when pure, is recognized by the fact that a splinter of wood, glowing at one end, bursts into flame when introduced into the gas. Only one other gas (see nitrous oxide) behaves similarly. The Measurement of Combining Proportions. In a number of condensed statements we have given the proportions by weight of the materials combining. It is now desirable that we should know how the necessary measurements are made. The most exact measurement of the proportions in which the elements AIR AND OXYGEN 37 combine to form compounds involves manipulations too elaborate to be gone into here. One or two brief statements, diagrammatic rather than accurate, will show the principles, however. If we take a weighed quantity of iron in a test-tube and heat it with more than enough sulphur (an excess of sulphur), we get free sulphur along with the ferrous sulphide (p. 14), and no free iron survives. We may remove the free sulphur 'HUH '. ^" - by washing the solid with carbon disulphide. The difference between the weights of the ferrous sul- phide and the iron gives the amount of sulphur combined with the known FlG 21 quantity of the latter. As an example of the study of the combination of a metal with oxygen, we may weigh a small amount of copper in the form of powder in a porcelain boat and pass oxygen over the heated metal (Fig. 21). If we limit the oxygen, part of the copper may remain unaltered; if we use it freely, the excess will pass on unchanged. The original weight of the copper, and the increase in weight, representing oxygen, give us the data for determining the compo- sition of cupric oxide. The data furnished by one rough lecture- experiment, for example, were as follows: Weight of boat -f- copper ~" 4.278g. Weight of boat empty 3 . 428 g. Difference = weight of copper . 860 g. Weight after addition of oxygen 4.488g. Weight without oxygen 4 . 278 g. Difference = weight of oxygen 0.210 g. The proportion of copper to oxygen, so far as this one measure- ment goes, is therefore 85 : 21. The results of quantitative experiments are often recorded in the 38 TEXT BOOK OF CHEMISTRY form of parts in one hundred. To find the percentage of each con- stituent, we observe that the proportion of copper is 85 : 85 + 21, or T 8 o% of the whole. That of the oxygen is T 2 oV of the whole. Thus the percentages are : Copper, 106 : 85 : : 100 : x. x = 80.2. Oxygen, 106 : 21 : : 100 : x'. x' = 19.8. Naturally, the mean oT the results of a number of more carefully managed experiments will be nearer the true proportion. The per^ centages at present accepted as most accurate are 79.9 and 20.1. In the case of mercuric oxide, we may decompose a known weight of the oxide (p. 15), collect the mercury and weigh it, and ascer- tain the oxygen by difference. The names of the constituent elements in a compound, together with the proportion by weight in which they are present, are called the composition of the substance. Thus, the composition of cupric oxide is copper : oxygen : : 79.9 : 20.1. This is the percentage com- position, but other numbers expressing the same proportion (such as 63.57 : 16) will serve the purpose. All experiments involving measurement, such as those used in determining composition, are called quantitative experiments. Another Quantitative Experiment. The following will show how the combining proportions may be measured when the product is a gas, the weight of which must be ascertained. Sulphur burns in oxygen to form sulphur dioxide. A known weight of sulphur is placed FlG - 22 - in a porcelain boat (Fig. 22) , which has already been weighed. The U-shaped tube to the right contains a solution of potassium hydroxide, which is capable of absorbing the resulting gas. The oxygen enters from the left. When the sul- phur is heated, it burns in the oxygen, and the loss in weight which the boat undergoes shows the amount of sulphur consumed. AIR AND OXYGEN 39 The gain in weight of the U-tube shows the weight of the compound produced. By subtracting, we get the quantity of oxygen. In one experiment, the loss in weight of the boat and its con- tents (= sulphur) was 1.21 g. The weight gained by the U-tube was 2.42 g. The difference (= oxygen) is 1.21. The proportion of sulphur to oxygen in sulphur dioxide is therefore 1.21 : 1.21 or 1 : 1 or, in percentages, 50 : 50. This proportion is very close to the accepted value, 32.06 : 32. The same method could be used for carbon, for the carbon dioxide produced would be absorbed in the solution of potassium hydroxide. Combustion. Violent union with oxygen is called, in popular language, combustion or burning. Yet, since oxygen is only one of many gaseous substances known to the chemist, and similar vigorous interactions with these gases are common, the term has no scientific significance. Even the union of iron and sul- phur gives out light and heat, and is quite similar in the chemical point of view to combustion. A misleading term often used in this connection is kindling temperature. It gives the impression that there is a definite tem- perature at which combustion will start. But the temperature is only one of the conditions which produce combustion. Finely powdered iron will start burning at a lower temperature than will an iron wire, because it presents relatively more surface to the gas. Again, if the oxygen is at less than one atmosphere pressure, the wire will require to reach a higher temperature before combustion will begin. Finally, the vapor of methyl alcohol and air requires to be raised above a red heat before combustion starts, but a pocket cigar-lighter sets fire to this very mixture by means of a contact agent (a thin platinum wire) without any other means of heating being required. So that, the conditions under which combustion begins involve the physical condition of the solid, the pressure of the gas or vapor, the presence or absence of a contact agent and 40 TEXT BOOK OF CHEMISTRY the nature of the contact agent, as well as the temperature. No definite kindling temperature can be given, unless the other con- ditions are specified also. Kindling conditions involve several variables, of which the temperature is only one. Oxidation. The slower union with oxygen which occurs in rusting is called oxidation. We shall see later, however, that it has been found convenient to stretch this term so as to cover com- binations of other elements than oxygen, and even to include actions not involving combination. At this point we can discuss only oxidation by oxygen. This process of slow oxidation by oxygen, although less con- spicuous than combustion, is really of greater interest. Thus the decay of wood is simply a process of oxidation whereby the same products are formed as by the more rapid ordinary combustion. Sewage is mixed with large volumes of river water, the object being, not simply to dilute the sewage, but to mix it with water containing oxygen in solution. This has an oxidizing power like that of oxy- gen gas and, through the agency of bacteria, quickly renders dis- solved organic matters innocuous by converting them for the most part into carbon dioxide and water. Thus, a few miles further down the stream, the water becomes as suitable for drink- ing as it was before the sewage entered. In our own bodies we have likewise a familiar illustration of slow oxidation. Avoiding details, it is sufficient to say that the oxygen, from the air taken into the lungs, combines with the haemoglobin in the red blood- corpuscles. In this form of loose combination, it is carried by the blood throughout our tissues and there oxidizes the foodstuffs which have been absorbed during digestion. The material prod- ucts are carbon dioxide and water, of which the former is carried back to the lungs by the blood, and finally reaches the air during exhalation. The important product, however, is the heat, given out by the oxidation, which keeps the body warm. The opposite of oxidation, the removal of oxygen, is spoken of AIR AND OXYGEN 41 in chemistry as reduction. But this term, also, has been stretched to cover other kinds of chemical change. Spontaneous Combustion. Sometimes a mere slow oxi- dation develops into a combustion, which is then known as spon- taneous combustion. To understand this, we must note the fact that a given weight of material, say, iron, in combining with oxy- gen to form a given oxide, will liberate the same total amount of heat whether the union proceeds rapidly or slowly. If the action proceeds slowly, and the material being oxidized is freely exposed to the air, the latter will become heated and will carry off the heat as fast as it is produced. Thus, no particular rise in temperature will occur. If, however, the material is a poor conductor of heat, like hay or rags, and there is sufficient air for oxidation, but not enough to carry off the heated air, the heat may accumulate and a temperature sufficient to start combustion may be reached. Such a situation sometimes arises in hay-stacks. It occurs also when rags, saturated with oils used in making paints (linseed oil and turpentine) are left in a heap. These oils, in " drying," combine with oxygen from the air and turn into a tough, resinous material. The rags, being poor conductors of heat, finally become hot enough to burst into flame, and serious conflagrations often owe their origin to causes such as this. Oily rags should always be disposed of by burning, or should at least be placed in a closed can of metal. Fires in coal bunkers of ships arise from the same cause slow oxidation, with accumulation of the resulting heat. That coal does undergo slow oxidation, especially when freshly mined, is shown by the fact that such coal, if left exposed to the air for months, may lose from 2 to 5 per cent or more of its heating value. Uses of Oxygen. A number of the practical applications of oxygen have already been mentioned. For example, in the fore- going section we have referred to its use in breathing, its role in decay, which is a beneficent process because it removes much 42 TEXT BOOK OF CHEMISTRY useless matter which might otherwise cause disease, and its value in the disposal of sewage. Power and heat for commercial pur- poses are almost all obtained by the burning of coal, in which oxy- gen from the air plays a large part. If we had to purchase the oxygen as well as the coal, we should require at least three tons of oxygen for every ton of coal. Oxygen in cylinders and oxygen generators are used to restore the supply in the atmosphere of submarine boats, as well as for the purposes already mentioned (p. 41). Substances Indifferent to Oxygen. Finally, since the atmosphere contains so large a proportion of oxygen, substances which do not oxidize and, when heated, do not burn, have many uses. Gold, silver, and platinum are of this kind (p. 36), and are used for ornaments. The last is used for crucibles in which bodies are heated in the laboratory. Although iron burns in pure oxygen, it does not oxidize rapidly in the air even when heated, and so is used for making vessels for cooking and in constructing fireproof buildings. Compounds, already fully oxidized, are naturally not com- bustible. Of this nature are sandstone, granite, brick, porcelain, glass, and water. All these are, therefore, fireproof. More- over these substances do not give off oxygen when heated (water decomposes slightly). Glass and porcelain thus neither lose nor gain in weight when heated, and are suitable materials for labora- tory apparatus. Exercises. 1. What percentage by weight of free oxygen is obtained by heating: (a) mercuric oxide, (6) potassium nitrate, (c) potassium chlorate? At $1.50 (7/8), $0.15 (8d), and $0.15 (8d) per kilogram, respectively, which is the cheapest source of oxygen ? 2. Using the data on p. 34, calculate the weight of oxygen dis- solved by 100 c.c. (= 100 g.) of water at 20. 3. Why does a forced draft make a fire burn more rapidly? AIR AND OXYGEN 43 4. Why does a naked flame sometimes cause an explosion in a mine, when the air of the mine is filled with coal dust ? 5. The substances, like phosphorus and sulphur, which burn rapidly in ordinary oxygen, combine very, very slowly with oxygen which has been freed from moisture by careful drying. How is this effect of water to be classified ? 6. Air is 20 per cent oxygen. Why does iron burn brilliantly in pure oxygen, but not in air? 7 What weight of copper would be required to combine with the amount of oxygen contained in 100 grams of mercuric oxide? CHAPTER IV THE MEASUREMENT OF GASES. THE MOLECULAR HYPOTHESIS AFTER the discussion in regard to the proportions of oxygen in the air and the measurement of the volume of oxygen removed (Fig. 15, p. 26), it will readily be imagined that measuring the volume of a sample of gas is a common operation in chemistry. Indeed, it is much easier to measure quantities of gas by noting their volumes than by weighing them. Some facts have to be taken account of, however, in order that the measurements of the volume may be of value. A sample of gas diminishes in volume when the pressure in- creases, and it increases in volume when the temperature rises'. Hence, when the volume of the gas is measured, the pressure and the temperature must also be stated. FIG. 23. Measurement of Pressure. In order that the pressure may be easily measured, we arrange the sample of gas so that its pressure is the same as that of the atmosphere at the moment. To do this, if, for example, the gas is contained in the narrow tube (Fig. 23), we bring the water inside the tube to the same level as that outside by lowering the tube. Then we read the volume by means of the graduation (not shown) on the tube and at the same time we read the height of the barometer. Correction of the Volume to 760 mm. Pressure. Since the barometer varies in height, and the volume of a sample of gas therefore varies also, it is convenient to " correct" the volume by (( reducing" it to that which the gas would occupy at "standard" 44 THE MEASUREMENT OF GASES 45 pressure, namely 760 mm. Now, the volume of a sample of gas varies inversely with the pressure (Boyle's law). If, for example, the volume is 23 c.c. and the observed pressure of the barometer is 745 mm., Boyle's law enables us to calculate the volume the same sample of gas would occupy at 760 mm. The volume changes in the ratio of these two pressures. If the pressure of the gas were actually changed to 760 mm., a greater pressure the gas would assume a smaller volume. Hence, the new volume = 23 X = 22.5 c.c. That is to say, if the new volume is to 7oO be less, we place the smaller pressure in the numerator. If a sample of gas occupies 15 c.c. at 850 mm., what volume will it occupy at 500 mm. ? Here the new pressure is smaller, and the 850 new volume therefore greater. New volume = 15 X = 25.5 c.c. Correction of the Volume to C. All gases at are found to gain 1/273 of their volume when heated 1 degree, 2/273 for 2 degrees and 273/273 for 273 degrees. Thus at 273 the volume is doubled. When cooled below 0, every gas similarly loses 1/273 of its volume for each degree. At 273, if the regular contraction continued so far, the sample would, by cal- culation, at least, lose all its volume. In point of fact, however, all gases liquefy before the temperature has fallen to 273. The rule contained in these statements is known as Charles' law. By applying an arithmetical device, we can state the law in a form which makes its use in calculations quite easy. The device consists in adding 273 algebraically to all temperatures. The temperature, when 273 has been added, is called the absolute temperature. The rule then reads: The volume of a sample of any gas is directly proportional to the absolute temperature. Thus, a sample of gas occupies 45 c.c. at 15, what would be its volume at 10? After we have applied the device, this reads: a sample of gas occupies 45 c.c. at 15 -f 273 = 288 Abs., what 46 TEXT BOOK OF CHEMISTRY would be its volume at 10 + 273 = 283 Abs. ? The volume changes in the ratio of these absolute temperatures. Since the new temperature (10 C. or 283 Abs.) is lower, the volume be- comes smaller. Therefore, putting the smaller number in the numerator, the volume at 283 = 45 X -- = 44.2 c.c. 288 Again, a sample of gas occupies 125 c.c. at 25, what will be its volume at 15 ? The absolute temperatures are 25 + 273 = 298 Abs., and 15 + 273 = 258 Abs. As the new temperature is lower, the volume will be less. Hence the new volume 125 X = 108.2 c.c. In practice, when a sample of gas is measured, we read the existing temperature, and correct the volume to that which the sample specimen would occupy at C. For example, the volume is 102 c.c. at 18, what is it at 0? The absolute temperatures are 18 + 273 = 291 Abs., and + 273 = 273 Abs. The new 273 volume will be smaller. Hence, the new volume is 102 X = 95 c.c. This correction enables us to compare all volumes of gases as if they had been measured at the same standard temperature, namely at 0. Corrections for Pressure and Temperature Combined. Since the volume changes, due to alterations in pressure and in temperature, are independent of one another, the corrections may be made either separately or together. The latter is more convenient. Thus, a sample of gas occupies 190 c.c. at 17 and 750 mm., what will be the volume under standard conditions, namely and 760 mm. pressure? 970 7 en New volume = 190 X X = 176.6 c.c. Correction for the Tension of Aqueous Vapor. When a gas is measured over mercury, the latter gives off practically THE MOLECULAR HYPOTHESIS 47 no vapor at room temperature, and the foregoing are the only corrections required. If, however, the sample of gas is stand- ing over water, then the volume is not that of the gas, but that of the gas plus a certain amount of water vapor. The latter must be subtracted. To do this we have to remember that, in a gaseous mixture, each one of the several gases (or vapors) exercises the same pressure as if it were present alone (Dalton's law of partial pressures). Now the pressure of the water vapor, in a gas stand- ing over water, at each temperature is known (see p. 62 and Appendix IV). It varies from 13.5 mm. at 16 to 23.5 mm. at 25. When, therefore, the gas is measured over water, the pres- sure of the water vapor is subtracted from the barometric reading, before the above-mentioned corrections are applied. For example, a specimen of a gas, standing over water, occupies 175 c.c. at 19 and 752 mm., what is the volume of the same gas at and 760 mm. when the gas is free from water? At 19 the aqueous tension is 16 mm. The pressure of the gas alone is therefore 752 16 = 736 mm. The fully corrected volume = Molecular Hypothesis. The relations between pressure and volume (Boyle's law) and between either of these and temperature (Charles' law) in gases is explained by the molecular hypothesis. According to this idea, all matter is composed of minute particles called molecules, those of any given substance being all alike in nature and in mass. In solids and liquids these molecules are closely packed to- gether. In gases, however, they are widely scattered, with much vacant space between them. A gas is in fact a vacuum, with numerous relatively minute particles scattered through it. When a gas is compressed, only the spaces between the molecules are reduced. By assuming further that, in gases, the molecules are in rapid motion, and produce pressure by striking the walls 48 TEXT BOOK OF CHEMISTRY of the vessel, and that this motion is increased by raising the temperature, all the laws of gases can be completely ex- plained. This hypothesis was first stated in rather complete form by Lomonossov (1744-50), and later, independently, by Water- ston (1845) and by Clausius (1857). Exercises. 1. Show that, if the levels of the water inside and outside the tube (Fig. 15, p. 26) are equal, the pressure of the gas inside must be equal to the atmospheric pressure. 2. Find the volume that 48 c.c. of gas at 732 mm. would occupy at 760 mm. 3. Reduce 48 c.c. of gas at 780 mm. to standard pressure (760 mm.). 4. Find the volume which 28 c.c. of gas at 775 mm. would occupy if the pressure changed to 730 mm. 5. Find the volume which 320 c.c. of gas at 20 would occupy at (pressure unaltered) . 6. Reduce 600 c.c. at 25 and 760 mm. to standard conditions (0 and 760 mm.). 7. What change in volume would occur if 1 liter of gas at 200 were cooled to 0. 8. Reduce 260 c.c. of gas at 10 and 742 mm. to standard conditions. 9. Reduce 35 c.c. of gas standing over water at 21 and 732 mm. to dry gas under standard conditions. 10. Give the feature of the molecular hypothesis which corre- sponds to each of the following facts about gases: (a) A specimen of gas can be compressed to 1 /2000th of the volume it occupies at 760 mm. (b) A gas, even if heavier than air, gradually leaves an open cylinder. (c) A specimen of a gas exercises pressure equally in all ' directions. THE MOLECULAR HYPOTHESIS 49 (d) Crowding the specimen of gas into one-third of its original volume triples its pressure. (e) When the temperature is raised, the pressure of a speci- men of gas increases (if the volume remains un- changed). 1 1 . When we drink lemonade through a straw, what causes the liquid to flow up into our mouths? CHAPTER V HYDROGEN AFTER considering the atmosphere, and particularly its most active component, we naturally turn to water, which, like air, is so closely associated with our daily life. We find that water is a compound of oxygen with hydrogen, and the latter element, therefore, next claims our attention. Hydrogen is of interest upon its own account. It is often used for filling balloons. Il- luminating gas, of the kind (water gas) used in most large cities, consists of hydrogen to the extent of about 40 per cent. Preparation by the Action of Metals on Water. Hydro- gen is not easily liberated from water, since this oxide of hydrogen, like many other oxides, is very stable. It is necessary to use some element with which oxygen will combine even more eagerly than with hydro- gen, and to offer this element in exchange for the hydrogen. If a piece of one of the very active metals, such as potassium, sodium, or calcium (now obtainable in commerce) is thrown into water, hydrogen is liberated and comes off in bubbles. The former two metals are lighter than water, and run about on the sur- face. The action with potassium is so violent that the hydrogen usually catches fire, and both with sodium and potassium much heat is produced. The action often ends with a slight explosion, so that a glass plate should be held up to protect the eyes. This experiment is too dangerous to be tried by a novice. With calcium the action is rapid, but not violent, and there is no danger. The metal sinks to the bottom of the vessel (Fig. 24), so that a 50 HYDROGEN 51 test-tube or bottle filled with water can be inverted over the metal to catch the gas as it ascends. The metals, of course, act upon a small part only of the whole vesselful of water. In each case the metal displaces one-half of the hydrogen from that part of the water upon which it acts: Sodium (23) -f- Water (18) - Hydrogen (1) + Sodium hydroxide (40). Hydrogen Sodium Hydrogen Hydrogen Oxygen Oxygen The products are hydrogen, along with potassium hydroxide, sodium hydroxide, or calcium hydroxide. In each case the name of the hydroxide indicates the composition, as including the metal along with hydrogen and oxygen. The first two hydroxides are very soluble in water, but most of the calcium hydroxide is not dissolved, and may be seen suspended in the liquid. Magnesium will liberate hydrogen from water, provided the latter is hot. If steam be passed through a heated tube con- taining iron filings, a mixture of hydrogen, with unused steam, issues at the other end. The magnetic oxide of iron, not hydroxide, remains in the tube: Iron + Water > Hydrogen + Magnetic oxide of iron. Hydrogen Iron Oxygen Oxygen This is a method much used in making hydrogen for commercial purposes. Silver, gold and platinum, which do not combine directly with pure oxygen, and even copper and mercury, which do, are all unable to form oxides and to liberate hydrogen when heated in steam. Displacement. The foregoing actions, in which hydro- gen is liberated, present us with a new the third variety of chemical change. Here an element displaces one of the elements from a compound, setting it free, and unites with the rest of the 52 TEXT BOOK OF CHEMISTRY constituents of the compound. Thus, calcium displaces part of the hydrogen, and unites with the oxygen and the rest of the hydrogen. Preparation by the Action of Metals upon Dilute Acids. All the metals which displace hydrogen from water or steam will also act upon cold dilute acids and displace the hydrogen they contain. This action is the one most commonly employed in the laboratory (Fig. 25). The gas, being much lighter than air, is collected by downward dis- placement of the air (Fig. 26b). Heavy gases are col- lected by upward displacement of air (Fig. 26a). Extremely pure zinc is almost inactive, but com- mercial zinc, on account of the contact action of the slight impurities it contains, gives a steady, not too violent, evolution of hydrogen. Sulphuric acid and hydrochloric acid, in each case diluted with water, are convenient acids. Iron shows about the same activity, but, on account of the impurities FIG. 25. FIG. 26a FIG. 26b. usually present in iron filings or wire, the hydrogen contains other gases which exhibit a distinct odor. Zinc (65.4) + Sulphuric acid (98) -> Hydrogen (2) -f Zinc Sulphate (161.4). ( Qnlv^Vm*. t O.-.lttU.-.. Hydrogen * v/iygeii Ferrous chloride (126.9) 1 Oxygen Iron (56) -f Hydrochloric acid (72.9) Hydrogen (2) Hydrogen Chlorine Iron Chlorine It will be seen that the action is of the form we have called displacement (p. 51). HYDROGEN 53 The Proportions by Weight. It may be well to remind ourselves that the weights of the various materials (given in brackets) are obtained by laboratory measurements. It is always found that the total weight of the product is exactly equal to that of the materials used (p. 22). Also, that a given weight of the metal, say zinc, will always displace and liberate the same weight of hydrogen. Also, that the proportions by weight of the con- stituent elements in the compound produced are always the same (p. 22). If we place a weighed piece of zinc in hydrochloric acid, and wait until the zinc has all disappeared, we can then boil away the water and unused acid, and weigh the white, solid zinc chloride. We find that 65.4 grams of zinc always leave 136.3 grams of zinc chloride. The difference, 70.9, is the chlorine, and whatever weights we take, the proportion of zinc to chlorine in the zinc chloride is always in the ratio 65.4 : 70.9. Zinc (65.4) + Hydrochloric acid (72.9) -> Hydrogen (2) + Hydrogen 2 Zinc 65.4 Chlorine 70.9 Chlorine 70.9 Chemically Equivalent Quantities. It will be observed that 65.4 parts of zinc displace 2 parts of hydrogen, whether the acid used is sulphuric acid or hydrochloric acid. The propor- tion is, in fact, the same with every acid. Hence 65.4 parts of zinc and 2 parts of hydrogen are spoken of as chemically equi- valent quantities. The quantities of the displacing and of the displace delement are in all cases referred to as chemically equiv- alent. The term equivalents is applied also to the quantities liberated by decomposition of a compound, like the 200 parts of mercury and the 8 parts of oxygen (p. 19). It is likewise used of the pro- portions combining when chemical union takes place, as in the case of phosphorus 31 parts and oxygen 40 parts (p. 35). The proportions of the element in zinc chloride (p. 53) are also chemi- cally equivalent. 54 TEXT BOOK OF CHEMISTRY Chemically equivalent quantities (or, simply, equivalents) of two substances are exact quantities which enter into or result from a chemical reaction. The Order of Activity of the Metals. It will greatly aid us in remembering a number of the facts already given, as well as many others, if we compare carefully with those facts the order in which the metals stand in the adjoining list. The most active metals are at the top. All above hydrogen displace this element from dilute acids (and, with more difficulty, from water); those below it do not. The first displaces the hydrogen from water so violently that the gas catches fire, the second dis- places it vigorously, the third less vigorously. Mag- nesium requires hot water and iron superheated steam. Copper and the metals following it do not liberate hydrogen from water. Again, the upper metals act too violently on dilute acids, and zinc is used to prepare the gas. Copper and the metals following it do not displace hydrogen from dilute acids. Still again, we recall the fact (p. 36) that, when we heat metals in pure oxygen, the last three do not become oxidized at all. Those preceding silver do t combine with pure oxygen mercury with difficulty, and the others mom fcxiti more vigorously as we ascend the list. On ' i?ofefter hand, if we start with the oxides of all the metals, we find that those at the foot of the list, up to and including mer- curic oxide, lose all their oxygen when heated, leaving the metal. Other facts of a similar nature will be mentioned as we en- counter them. Meantime, it may be noted that the metals found uncombined in nature are those following hydrogen. Again, ORDER OF ACTIVITY. Metals Potassium Sodium Calcium Magnesium Aluminium Manganese Zinc Chromium Iron Nickel Lead Tin Hydrogen Copper Bismuth Antimony Mercury Silver Platinum Gold HYDROGEN 55 the metals known to have been first used by the human race were gold and silver. In the " bronze age" means of liberating copper from its ores had been discovered. Lead, tin, and iron came later. The list, read from the bottom up, gives, therefore, roughly, the historical order in which the metals came into use. Preparation by Electrolysis of a Dilute Acid. A con- venient way of obtaining pure hydrogen is by passing a current of electricity through a dilute add (Fig. 27). The gas is liberated at the negative wire (cathode) and collects in the tube (also filled with the dilute acid). The direct, 110-volt current, passing through a 16-c.-p. lamp placed in series with the electro- lytic apparatus, may liberate 50 c.c. of hydro- gen in 7-8 minutes. Every acid contains hydrogen, combined with other elements. The other elements are carried to the positive plate (anode) and therefore do not interfere with the collection of pure hydrogen. What may be liberated at the positive plate depends upon the acid used. With hydrochloric acid, it is chlorine; with sulphuric acid, oxygen comes off and sulphuric acid is regenerated. Hydrochloric acid (in Aq. Soln.) > Hydrogen (neg. plate) -f- Chlorine (pos. plate). The process is called electrolysis, from the Greek, meaning decomposed by electricity. Physical Properties of Hydrogen. The gas is colorless, odorless, and tasteless. It is, bulk for bulk, the lightest known gas, the density of ah- being about 14.5 times as great. It can be liquefied by compression below 234 (its critical temperature). It dissolves in water to the extent of 1.8 volumes in 100 volumes of water at 15. 56 TEXT BOOK OF CHEMISTRY The lightness of the gas may be shown by pouring it upwards from one jar to another, or by balancing an inverted beaker with shot, and allowing hydrogen to flow in and displace the air. Several metals can absorb (" occlude") hydrogen gas. Iron FIG. 28. FIG. 29. takes up about 19 times its own volume, platinum 50 volumes, and palladium from 500 to 800 volumes. Liquid hydrogen, when allowed to evaporate rapidly, freezes to a colorless solid, which melts again at 260. Chemical Properties of Hydrogen. That hy- drogen, when it burns in the air, forms water, was first shown by Cavendish (1781). A test-tube or beaker, filled with cold water and held over a flame of. burning hydrogen, will condense the steam to droplets of water (Fig. 28). Hydrogen + Oxygen Water. The union is very violent, so that when a mixture of hydrogen with pure oxygen is set on fire (they do not unite when cold) much heat is liberated in the explosion. The gases can be made to burn quietly, but with an exceedingly hot flame, by the use of an oxy-hydrogen burner (Fig. 29), which is constructed like a blast lamp. Iron melts and burns in the flame. A piece of quicklime, held in the FIG. 30. HYDROGEN 57 flame, glows with a brilliant white light the calcium light (Drummond light or lime-light). For such uses the gases are obtained in compressed form in iron cylinders (Fig. 30). Hydrogen combines vigorously with chlorine, giving hydrogen chloride, a gas of which hydrochloric acid is a solution. It unites with the three most active metals in the list (p. 54). Calcium hydride is sold under the name of hydrolyte, and is used, on ac- count of its action on water, as a source of hydrogen. Hydrogen acts upon many compounds containing oxygen, re- moving the latter to form water. Thus when the oxide of iron or of copper is heated in a tube in a stream of hydrogen, water is produced and the metal remains : Magnetic oxide of iron + Hydrogen > Iron + Water. Iron Hydrogen Oxygen Oxygen Oxides of metals above iron in the "order of activity" (p. 54), however, are very stable. Hydrogen is unable to remove the oxygen from such oxides and leave the metal. Reduction. The removal of oxygen from a compound by its union with some other substance is called reduction and the substance (in the foregoing instance, hydrogen) is called a reduc- ing agent. Carbon, in the form of coal or coke, is the agent of this kind most commonly used in chemical industries. The term re- duction is applied to some other chemical actions, in which oxygen is not concerned. In all cases, however, reduction is the opposite of oxidation (p. 35). Exercises. 1. Name three varieties of chemical change (pp. 9, 16, 51) and explain the difference between them. 2. What do you infer as to the composition of a substance when it is named: (a) an oxide, (b) an hydroxide (p. 51)? 3. What are the equivalent quantities of: (a) carbon and oxy- gen (p. 35), (b) zinc and sulphur? 58 TEXT BOOK OF CHEMISTRY 4. What law shows that the ratio of chemically equivalent quantities of any two substances must be constant? 5. Name the metals which: (a) do not liberate hydrogen from water of dilute acids, (b) are found free in nature. 6. Why does hydrogen gas, when poured out, flow upwards? Why is it an excellent gas for filling balloons ? CHAPTER VI WATER IF oxygen is necessary to the life of plants and animals, so also is water. The human body is saturated with it, and water to make up for evaporation, as well as to aid in digestion, is a most necessary part of our food. The ocean covers about three-fourths of the surface of the earth, and the "dry" land would be unin- habitable if it were really dry. The air always contains more or less water vapor. Measurement of the Composition by Weight. An ar- rangement by which the proportion by weight of hydrogen and FIG. 31. oxygen in water can be determined is shown in Fig. 31. The bulb A contains cupric oxide, which is heated. Hydrogen from a generator or cylinder enters on the left and reduces the oxide, forming copper and water: Hydrogen -f Cupric oxide > Water + Copper. 59 60 TEXT BOOK OF CHEMISTRY The water is carried as vapor by the excess of hydrogen and passes into the U-tube B. This tube contains calcium chloride, a sub- stance which absorbs water greedily and is used therefore for drying gases. Here the water is all caught, while the hydrogen passes on. The tubes A and B with their contents are weighed just before, and again just after, the experiment. The loss of weight in A is the weight of the oxygen. The gain in weight in B is the water. The difference be- tween these numbers is the hydrogen. It is found that the weights of hydrogen and oxygen thus ascertained always stand in a ratio close to 1 (Hyd.) :7.94 (Ox.) or 1.008:8, the proportions accurately determined by Morley. Measurement of the Composition by Vol- ume. The proportions by volume in which hydrogen and oxygen combine may be shown by introducing the two gases into a tube, filled with mercury and inverted in a cylinder of mercury (Fig. 32). The volumes, at atmospheric pressure, are read by lowering the tube, after the introduc- tion of each gas, until the levels of the mercury inside and outside are alike. A spark from an in- duction coil passed between the platinum wires inserted at the top of the tube causes the union of the gases. The water condenses to a slight dew and the volume of the single gas which remains is Thus, if 19.5 c.c. of oxygen and 20 c.c. of hydrogen are taken, the volume of gas remaining is 9.5 c.c. and this gas is afterwards found to be oxygen (test, p. 36). The volumes con- sumed were therefore 19.5 9.5 = 10 c.c. of oxygen and 20 c.c. of hydrogen. The ratio by volume is therefore 2 Hyd. : 1 Ox. If these exact proportions are used, the mercury fills the tube after the explosion, but is apt to break it by striking the top violently. FIG. 32. measured. WATER 61 By taking the gases in the exact ratio 1 : 2, and surround- ing the tube by a wider one through which steam passes, the condensation of the resulting steam is prevented. It is found that, when all the gases are measured at the same temperature (here about 100), a shrinkage of one-third occurs. That is to say: 1 vol. Ox. -f- 2 vols. Hyd. -> 2 vols. Steam. Gay-Lussac 9 s Law of Volumes. When other chemical actions between gases are studied in the same way, it is found that, in every case, the volumes of the gases used and produced in a chemical change can always be represented by the ratio of small whole numbers. This fact is exceedingly interesting. It was first discovered by Gay-Lussac in 1808. There is no such simple relation amongst the proportions by weight, which usually can be expressed only by large numbers (see p. 36) or by irregular fractions, so that this evidence of the existence of a simple rule in regard to combining proportions is the first we have encountered and is very welcome. The use made of it by the chemist will be developed in the next chapter. Physical Properties. Water is without odor or taste. It is very pale blue in color, as is shown when we look through a con- siderable depth of water. Melting ice and freezing water have the same temperature. The point at which the mercury column of a thermometer stands, when the instrument is immersed in such a mixture of ice and water, is marked on the centigrade scale. This is the freezing- point of water. The density of ice is little over nine-tenths that of water. When water is heated, it gives off vapor more and more freely until finally it boils. This point is recognized by the fact that bubbles of vapor form within the liquid, rise through it, and burst on the surface. The temperature of the water (if it is pure), and of the 62 TEXT BOOK OF CHEMISTRY steam, are now found to be identical, and this thermometer read- ing is marked 100 C. the boiling-point. Water is thus the standard substance used in the graduating of thermometers. When so used for fixing the temperature of 100, the atmospheric pressure must be normal, that is 760 mm., for the temperature of boiling is lower when the pressure is lower. When water has reached the boiling-point, the temperature ceases to rise, and the heat supplied is used in changing the water into steam. The evaporation of 1 gram of water consumes 537 heat units or B calories (heat of vaporization). The melting of 1 gram of ice consumes 79 calories (heat of fusion). The calorie is the average amount of heat required to raise 1 gram of water one degree in temperature between 1 and 100. The amount of heat required to raise the tem- perature of a given mass of water one degree is greater than that required for an equal mass of any other common material. Hence, the temperature of the sea changes more slowly, and within a smaller range, than that of the rocks which compose the land. For this reason the climate of islands sur- rounded by much water is less variable from season to season within the year than is that of the continents. Steam. The quantity of vapor given off by water at various temperatures is best measured by the gaseous pressure it exer- cises. A little water introduced into a barometric vacuum (B, Fig. 33) will depress the mercury, and the difference in height is a measure of the vapor pressure of the water. The temperature may be changed by putting hot water into the tube surrounding the barometer, and thus the increase in vapor pressure with rising temperature may be shown. At the column of mercury is de- pressed 4.6 mm., at 10 the vapor pressure becomes 9.2 mm., FIG. 33. WATER 63 at 20 17.4 mm. (see Appendix IV). At 100 the level of the mercury in the tube would be depressed 760 mm., and would sink to the level of that in the trough. If a little air is first placed above dry mercury, causing it to fall, the additional depression pro- duced by adding water is the same as if the air had been absent (p. 47). Steam at 100, subjected to a pressure over one atmosphere, is condensed to water. Steam cannot be condensed by pressure, however, if the temperature is over 358, its critical temperature. At 100 the steam occupies more than 1700 times the volume of an equal weight of water. Steam is a perfectly invisible gas. The visible cloud of fog, issuing from a valve when steam escapes, is composed of min- ute drops of water formed by condensation. Molecular Relations of Liquid and Vapor. When the water was introduced above the barometric column, the vapor, or gaseous water, could have resulted only from the spontaneous motion of the molecules in the liquid. Some of the molecules, moving near the surface, went off into the space above the water and became gaseous. To be consistent, we must also conclude that the vapor above the water is not composed of the same set of molecules one minute as it was during the preceding minute. Their motions must cause many of them to plunge into the liquid, while others emerge and take their places. When the water is first introduced, there are no molecules of vapor in the space at all, so that emission from the water predominates. The pressure of the vapor increases as the concentration of the molecules of vapor becomes greater, hence the mercury column falls steadily. At the same time the number of gaseous molecules plunging into the water per second must increase in proportion to the degree to which they are crowded in the vapor. The rate at which mole- cules return to the water thus begins at zero, and increases steadily; the rate at which molecules leave the water maintains a constant 64 TEXT BOOK OF CHEMISTRY value. Hence the rate at which vapor molecules enter the water must eventually equal that at which other water molecules leave the liquid. At this point, occasion for visible changes ceases and the mercury comes to rest. We are bound to think, however, of the exchange as still going on, since nothing has occurred to stop it. The condition is not one of rest but of rapid and equal exchange. Such, described in terms of molecules, is the state of affairs which is characteristic of a condition of equilibrium. The condition is dynamic, and not static. Equilibrium. This term is used so often in chemistry, and is used in so unfamiliar a sense, that the reader should consider attentively what it implies. Three things are characteristic of a state of equilibrium: 1. There are always two opposing tendencies which, when equi- librium is reached, balance each other. In the foregoing instance, one of these is the hail of molecules leaving the liquid, which is constant throughout the experiment. It represents the vapor ten- sion of the liquid. The other is the hail of returning molecules, which, at first, increases steadily as the concentration of the vapor becomes greater. This is the vapor pressure of the vapor. These have the effect of opposing pressures and, when the latter becomes equal to the former, equilibrium is established. In all cases of equilibrium we shall symbolize the two opposing tendencies by two arrows, thus : Water (liq.) =; Water (vapor). 2. Although their effects thus neutralize each other at equilib- rium, both tendencies are still in full operation. In the case in point, the opposing hails of molecules are still at work, but neither can effect any visible change in the system. Equilibrium is a state, not of rest, but of balanced activities. 3 (and this the chief mark of equilibrium). A slight change in the conditions produces, never a great or sharp change, but always, and instantly, a corresponding small change in the state of the WATER 65 system. The change in the conditions accomplishes this by favoring or disfavoring one of the two opposing tendencies. Thus, for ex- ample, when the temperature of a liquid is raised, the kinetic energy of its molecules is increased, the rate at which they leave its surface becomes greater, the vapor tension increases and, hence, a greater concentration of vapor can be maintained. The system, therefore, quickly reaches a new state of equilibrium in which a higher vapor pressure exists. A heap of matter on a table is not in equilib- rium, because addition of more material produces no response until, when a very great quantity is added, the table breaks. But a body on the scales is in equilibrium, for the addition of the smallest particle produces a corresponding inclination of the beam. In the preceding illustration, the evaporating tendency was favored by a rise in temperature. As an example of a change in conditions disfavoring one tendency, take the case where the liquid is placed in an open, shallow vessel. Here the condensing tendency is markedly discouraged, for there is practically no return of the emitted molecules. Hence complete evaporation takes place. Ele- vation of the temperature hastens the process. A draft insures the total prevention of all returns, and has therefore the same effect. The two methods of assisting the displacement of an equilibrium, and particularly the second, in which the opposed process is weak- ened and the forward process triumphs solely on this account, should be noted carefully. They are applied with surprising effec- tiveness in the explanation of chemical phenomena. The States of Matter. Most substances are known in at least three different states, namely, a solid, a liquid, and a gaseous form. There is no magic about the number three, however. Very many substances are known in more states than three. Thus sulphur has a vapor state, two liquid states, and several different solid forms. There are even five different forms of ice, and it seems probable that all solids can exist in several different solid states. 66 TEXT BOOK OF CHEMISTRY With all substances the changes from a solid to a liquid, and from a liquid to a gaseous state, always take place at some def- inite temperature, similar to the melting-point of ice and the boiling-point of water. Water as a Solvent. Water has a remarkable power of dissolving many other substances, and is said, therefore, to have great power as a solvent. In the simplest cases, at least, the operation appears to be purely mechanical and this is classed therefore as a physical property. Natural Waters* Rain is the purest natural water. As it is formed by condensation of water vapor, and has been in contact with the atmosphere only, it contains only oxygen and other gases dissolved from the air, together with a little dust. Sea water contains the greatest amount of dissolved material, namely about 3.5 per cent. River and, especially, well waters contain materials in solution which have been dissolved from the soil and the rocks. These materials act chemically upon soap so that well waters are more or less hard, while rain water is soft. We have already learned (p. 40) that natural waters may also contain bacteria, which give rise to putrefaction and disease. Many river waters contain large amounts of clay and other insoluble substances suspended in them. The suspended matter can be seen, as it renders the water turbid, but the bacteria are invisible, and organic matter and bacteria may be present in waters which look per- fectly clear. Purification of Water. The suspended impurities, includ- ing the bacteria, may be removed by filtration. City waters are often filtered through extensive beds of gravel, but this treat- ment will not remove all bacteria. In many cases, small amounts of alum, or alum and lime, are added to the water, and the sus- pended matter is then allowed to settle, which it does very quickly, WATER 67 in large basins or reservoirs. By this coagulation method (p. 427), all but a few of the bacteria are removed. Sometimes the re- maining organisms are destroyed by adding a little bleaching powder before the water is distributed (p. 307). In the household, the simplest appliance is the Pasteur filter. It consists of a tube of unglazed porcelain, closed at one end (Fig. 34), through the pores of which the water is forced by its own pressure. The cylinder ("bougie") should be cleaned daily with a brush to remove the mud and the organisms which collect in its surface. All forms of filters must be cleaned at short in- tervals. If this is not done, the organisms multiply and soon the filter polutes the water instead of puri- fying it. Filtration does not remove dissolved matter, and therefore does not soften hard water (p. 13). For this latter purpose washing powders are used in the laundry (see p. 400). Ah 1 organisms can be killed by boiling unfiltered water, but the boiling should continue for at least 10 to 15 minutes to be effective. Pure water for chemical purposes is prepared by distillation (Fig. 35). Dissolved solids remain in the flask (or boiler). The steam is condensed by cold water circulating in the jacket of the condenser. Freshly distilled water contains only gases dissolved from the air. If kept in a vessel, however, such water quickly dissolves traces of glass or porcelain. The purest water is made by using a platinum tube in the condenser and a platinum bottle as the receiver. Chemical Properties: Stability. As we should infer from the vigor with which its constituents combine, water is a very stable substance. When steam is superheated, hardly a trace of decomposition occurs. Even when the temperature reaches FIG. 34. 68 TEXT BOOK OF CHEMISTRY 2500 (far above a white heat), only 1.8 per cent of the vapor is broken up into oxygen and hydrogen. When the steam cools, these elements recombine: Water <=* Hydrogen + Oxygen. Water is so frequently used simply to dissolve chemical sub- stances that we should note carefully the cases (see pp. 68, 69 FIG. 35. following), which are relatively few, in which it enters into chemi- cal action. Hydrates. Many substances unite with water to give compounds called hydrates. Thus if we take zinc sulphate (p. 52) and dissolve it in water and allow the excess of the latter to evapo- rate, the solid appears in long transparent crystals. When these are dried with blotting paper and heated in a test-tube, they give off WATER 69 a large amount of steam. The hydrate of zinc sulphate decom- poses and leaves anhydrous (Greek, deprived of water) zinc sul- phate. The latter, when once more moistened, changes back into the hydrate. Zinc sulphate -j- Water <= Hydrate of Zinc sulphate. Many common chemicals are in fact such hydrates. Thus com- mon blue-stone, used in gravity batteries, is a hydrate of cupric sulphate. When heated, it loses water and leaves the colorless, anhydrous cupric sulphate. These are cases of simple com- bination and decomposition. Efflorescence. Some hydrates are so unstable that the water passes off, even at room temperature, when the hydrate is left in an open vessel. Thus crystals of washing soda (hydrate of sodium carbonate) crumble to powder (effloresce) when not kept in a closed vessel. Sodium carbonate -f Water <= Hydrate of sodium carbonate. If a crystal of a hydrate like this is placed in the barometric vacuum (Fig. 33, p. 62) it is found to give a considerable vapor pressure of water, so that its tendency to decompose, when the vapor is allowed to escape, is easily understood. The measured pressure of water vapor coming from such hydrates at ordinary temperatures is found to be greater than the average pressure of water vapor in the atmosphere. On the other hand, when anhydrous cupric sulphate and zinc sulphate, obtainable from the hydrates by heating only, are spread out in the air, they return slowly to the hydrated con- dition. They combine with the moisture in the air. The vapor pressure of water in the air is greater than the vapor pressure of water which these hydrates give off at ordinary temperatures. Calcium chloride absorbs water vapor (p. 60) because of its ten- dency to form a hydrate. 70 TEXT BOOK OF CHEMISTRY Other Chemical Properties. The cases where water com- bines directly with oxides such as quicklime, is a familiar reaction, used to make plaster: CaO + H 2 > Ca(OH) 2 some oxides give acids, for example sulphur dioxide. SO 2 + H 2 > H 2 SO 3 . Where it acts chemically upon substances dissolved in it (p. 120) will be noted later. Reversible Chemical Actions. The contrary effects upon an unstable hydrate of leaving the bottle open or closed, referred to on p. 69, deserves a moment's notice. When understood, it explains many things in chemistry. The hydration and dehy- dration are opposite directions of the same chemical change, and the condensed statements of the actions were written with the double arrow to indicate this (p. 69). When the bottle is closed, the water vapor is unable to escape and recombines with the anhydrous particles as fast as other particles of the hydrate de- compose. A reversible action therefore can not complete itself, if the products of the action are kept together and not allowed to sepa- rate. On the other hand, a reversible action will go to completion, if one of the products escapes, as the water vapor does when the bottle is left open. This idea enables us to answer several in- teresting questions. For example, why does steam decompose to the extent of 1.8 per cent at 2500 degrees, but -not any further? All its parts are alike and are therefore equally capable of decomposing. The answer is, because neither the oxygen nor the hydrogen is removed, or can easily be removed, from the steam, and so the completion of the decomposition is prevented by continual recombination oj these gases. When a reversible action has come to a standstill, we say that equilibrium has been reached. This means that two opposing tendencies are neutralizing one another's effects. When will reversible actions go to completion? The products must be of such a nature that they separate easily. In practice WATER 71 this happens when one is a gas or vapor, like the water vapor coining from a hydrate, while the other is not. The settling of one product as a precipitate, while the other stays in solution, is, as we shall see (pp. 125, 126), another common way in which the separa- tion occurs. When iron and water are heated in a closed vessel, the hydrogen and oxide of iron which are produced (p. 51) react with one another to give back water and iron. Iron + Steam + Hydrogen + Magnetic oxide of iron Hydrogen Iron Oxygen Oxygen In a closed vessel we could never use this method for preparing any quantity of hydrogen. To prepare hydrogen by this action we leave the tube open, and let the steam sweep the hydrogen out. This separates it effectually from the oxide of iron, and prevents the reversal of the action. Devices depending on mechanical principles like this are con- tinually used in chemistry for securing easy methods of prepar- ing substances. The reader can now answer for himself the question why we are able to prepare oxygen by heating mercuric oxide (p. 15), in spite of the fact that the action is reversible (p. 28). Exercises. 1. If 50 c.c. of hydrogen and 37 c.c. of oxygen are exploded in a closed tube, which gas remains and what volume of it is left? 2. Why do not bubbles of steam ordinarily form in water and rise through it at temperatures below 100 ? 3. How many calories of heat would be required to change 5 grams of ice at into steam at 100. 4. How could one find out how much solid matter was dissolved in a sample of water ? 5. Define: filtration, distillation, efflorescence, chemical equilib- rium. 72 TEXT BOOK OF CHEMISTRY 6. Explain how water, even at room temperature, gradually dries up (p. 65). 7. Why does a strong wind hasten the evaporation of water (p. 65)? 8. Why does the lower surface of a sheet of ice on the pavement melt in the sun before the upper one ? CHAPTER VII CHEMICAL UNITS OF WEIGHT. FORMULAE As we have seen (p. 61), when the volumes occupied by sub- stances in the gaseous condition, rather than the weights, are taken as the basis of measurement, the combining proportions are simple and are expressible by small whole numbers (Gay-Lussac's law). This shows that there must be some relationship, connected with chemical combination, between the amounts of different sub- stances contained as gases in equal volumes. It suggests that we might do well to take such amounts (weights of equal volumes) as the standard or unit quantities for chemical purposes. Now this is precisely what the chemist has found it in practice most convenient to do, and the present chapter deals with the units of material based upon comparing equal volumes. Illustrations of Gay-Lussac 9 s Law. Let us first famil- iarize ourselves with the volume-measuring point of view in chemical actions. The following are a few observed facts, be- ginning with the union of hydrogen and oxygen already discussed (p. 61): (1) Hydrogen (2 vols.) + Oxygen (1 vol.) - Steam (2 vols.). (2) Hydrogen (1 vol.) + Chlorine (1 vol.) -> Hydrogen chloride (2 vols.). (3) Chlorine monoxide (2 vols.) > Chlorine (2 vols.) + Oxy- gen (1 vol.). (4) Mercuric oxide (not volatile) > Mercury (2 vols.) -j- Oxy- gen (1 vol.). (5) Phosphorus (1 vol.) + Oxygen (5 vols.) > Phosphorus pentoxide (1 vol.). (6) Zinc (at 1000, 2 vols.) + Sulphur (at 1000, 1 vol.) -* Zinc sulphide (not volatile). 73 74 TEXT BOOK OF CHEMISTRY It will be noted that in some cases, like (2), there is no change in the total volume. In others there is a shrinkage, as in (1) and (5). In still others, like (3), where chlorine monoxide decomposes, there is an increase in volume. In (5), in order that all the ma- terials may be gaseous, the whole experiment must be done at a very high temperature (and at some suitable pressure). In (4) the mercuric oxide itself does not become gaseous, but decomposes, so that its own relative volume cannot be given. In (6) the zinc and sulphur can be combined as vapors at 1000. The product (zinc sulphide) will not remain gaseous at any temperature at which its volume could be measured, however, and so its volume is not recorded. It must be kept constantly in mind that the law applies to volumes in the state of gas or vapor only. There is no rule about the proportions by volume required for the chemical combination of liquids and solids. One can read these illustrations in different ways. For ex- ample: (1) A given volume of steam is formed by union of the same volume of hydrogen with half as great a volume of oxygen. (4) Mercuric oxide, when decomposed by heating, gives two volumes of mercury vapor and one volume of oxygen in every three volumes of the escaping gases. (5) One volume of phos- phorus vapor, together with an equivalent quantity of oxygen, will give one volume of the vapor of phosphorus pentoxide all being measured at the same temperature. In fact, whenever two vaporizable substances are amongst the factors and products of a chemical change, their volumes thus are either equal, or are to one another in the ratio of whole numbers. The Standard or Unit Volume. The volumes in the fore- going paragraph are simply relative, and the statements are true of gaseous volumes of any actual dimensions (large or small), provided only they bear the proper relationship, such as 2:1, 1 : 1, or 1 : 5, in each case. An actual value has been chosen, CHEMICAL UNITS OF WEIGHT. FORMULAE 75 however, for the volume which is the standard or unit in chemistry. This is the volume occupied by 32 grams of oxygen, which is 22.4 liters at and 760 mm. This volume is equal to that of a cube about 11 inches in height (Fig. 36). At other temperatures and pressures this volume, in order to contain the same amount of material, alters its value, in accordance with the laws of Boyle and Charles (p. 45). The reason for selecting this particular volume (see p. 76) will be readily seen, so soon as we shall have pre- sented the actual weights of various mate- rials which fill it. 22 4 LITERS FIG. 36. The Weights Filling the Unit Vol- ume, 22.4 Liters. The following table contains a few sample substances, and gives the weight (in grams) of each which, in the gaseous condition, at and 760 mm., occu- pies the cube 22.4 liters. In the cases of compound substances, like water, we have given also the weights of the constituent ele- ments which together make up the total weight of the compound filling the unit volume: WEIGHTS OF GASES IN 22.4 LITERS (AT AND 760 MM.). Substance. Total wt. Wt. ox. Wt. hyd. Wt. chlor. Oxygen 32 32 Water 18.016 16 2.016 Hydrogen 2 016 2 016 Hydrogen chloride 36 468 1.008 35.46 Chlorine 70 92 70 92 Chlorine monoxide 86 92 16 70 92 The first column ("Total wt.") gives the number of grams of each substance occupying, as a gas, 22.4 liters at and 760 mm. This is the standard or unit weight of that substance. In the case of water, which is a liquid at room temperatures and pressures, a known volume of steam is weighed. The volume is reduced by rule to and 760 mm., and the weight of 22.4 liters is calculated from the reduced volume and the measured weight. 76 TEXT BOOK OF CHEMISTRY These standard or unit weights of substances are commonly called molecular weights (see p. 84). Unit or Standard Weights of the Elements. Let us now examine the weights of the constituent elements making up a cubeful of each substance, as shown in the last three columns of the above table. We must first be sure we understand what these numbers are. They are combining proportions, such as we have given on previous occasions (e.g., pp. 19, 15, 30). They are equiv- alents (p. 53). We can use them in our condensed form for representing chemical changes: Oxygen (16) + Hydrogen (2.016) '"-> Water (18.016). Hydrogen (1.008) + Chlorine (35.46) -> Hydrogen chloride (36.468). We observe at once that the weights in the oxygen column, or the chlorine column, for example, are not identical. There was no reason to expect that they would be alike, since different sub- stances differ in composition. But we do observe that the weights of any one element are all multiples of one of the numbers, either by unity or some other whole number. Thus, for hydrogen, the weights are : 1 .008, 2.016 and 2.016. The unit weight of hydrogen chloride contains one unit of hydrogen, those of water and of hydrogen itself two units each. This is a very surprising, natural fact, and, better still, one for which we instantly perceive a use. This fact greatly simplifies our task of finding some way of expressing the composi- tions of substances in a simple manner. The fact does not apply to a few compounds, only. If our table had included all the hundreds of compounds of chlorine (for example) which are capable of being converted into vapor, we should have found, indeed, many multiples of 35.46 larger than the two units (70.92) in chlor- ine monoxide, but no number smaller than 35.46 and none which was not a multiple of 35.46 by a whole number. Clearly we shall find it convenient to accept 35.46 as the unit or standard weight of CHEMICAL UNITS OF WEIGHT. FORMULA 77 chlorine for expressing the compositions of its chemical compounds. Its use will enable us to state the exact composition of any given compound by simply giving the whole number (1, 2, 3, etc.) by which the basal weight 35.46 is to be multiplied in the given case. Similarly, we take 1.008 as the unit weight of hydrogen and 16 as the unit weight of oxygen. By including volatile compounds of other elements in our investigation, we can similarly pick out the most convenient unit for each element. Those unit weights for elements are often called also combining or reacting weights, and still more frequently atomic weights (see p. 86). A complete list of their values for all the known elements is given in a table inside the rear cover of this book. These numbers will hereafter be in constant use. The following table shows a larger number of substances and gives their formulae: Substance. Molar Weight. Weights of Constituents in Molar Weight. Hydrogen. Chlorine. i i 1 PL, Carbon. H 2 O. This shows the unit quantities of each individual substance, but not necessarily the correct proportions in the particular action in question. Thus, H 2 : O 2 is equivalent to 2.016 : 32, whereas the correct proportion (p. 75) in this action is 2.016 : 16, or 2 X 2.016 : 32. Also, 2 contains enough oxygen for 2 H 2 0. The statement must therefore be adjusted ("balanced") to take ac- count of these two facts. The correct equation for the action is therefore : Balanced equation: 2 H 2 + O 2 > 2 H 2 0. The final statement is called an equation because the weights of each element on the two sides are equal, and the total weight on each side is the same. This is in accordance with the law of conserva- tion of mass (p. 20), which holds for all chemical changes. 82 TEXT BOOK OF CHEMISTRY Further Discussion of the Same Subject. At this point we have completed the explanation of the derivation of the units of weight, and of the symbols and formulae used in chemistry. We are now in a position to proceed with the application of these con- ceptions, as they are developed in the next chapter. This would be possible, at least so far as strict logic is concerned. But, in shaping our course so as to reach the results by the shortest route, we have omitted a number of interesting and useful ideas. These were not indispensable links in our reasoning, but, now that we have covered the essential steps, a consideration of these ideas will be of great value as a review of the subject and a means of master- ing it thoroughly. Formulae and their uses come up so constantly in every part of the science that we must not omit any means of securing a perfect understanding of them. The following sections are therefore devoted to restating in other ways certain parts of the same subject. Why was 22.4 Liters Chosen as the Unit Volume ? It was chosen of such dimensions that no element should receive a unit weight smaller than unity. Hydrogen is the element which enters in the smallest proportions into chemical combination. In 22.4 1. of some compounds there is as little as 1.008 g. of this element (see table, p. 75). Any size would have served the purpose equally well theoretically. But one much smaller than 22.4 1. would have held a weight of hydrogen chloride (e.g.) containing less than one gram of hydrogen. If 22.2 1. had been chosen, this would have contained exactly 1 g. of hydrogen (in combination with 35.2 g. of chlorine). This volume was, in fact, at one time, the standard. But its use reduces all the numbers in the same proportion, and the unit weight of oxygen (in water, for example), then becomes about 15.9. The chief advantage of the scale = 16 is that, on that scale, very many of the atomic weights happen to be so close to whole numbers (see table) that the nearest whole number can be used for most purposes. This simplifies calculation. CHEMICAL UNITS OF WEIGHT. FORMULA 83 The accepted scale is, therefore, that of 32 for the unit weight of oxygen gas; 22.4 1. (the volume of 32 g. of free oxygen) for the chemical unit of volume; and 16 for the unit weight of oxygen in compounds containing that element. The unit weights of all the other compounds and elements are based upon this scale. Application of the Molecular Hypothesis in Chemistry Avogadro's Hypothesis. We have seen (p. 47) that the physical behavior of matter, and particularly of gases, is ex- plained by the conception that matter is composed of molecules. The easy compressibility of gases is explained by the smallness of the molecules and relative vastness of the empty space between them. Diffusion is explained by the continual motion of the molecules, and pressure by the innumerable blows which the molecules deliver when they strike the boundary walls of the space in which the gas is confined. Now, when gases interact chemically, the volumes required for complete interaction are either equal or stand in the ratio of small whole numbers. This law (Gay-Lus- sac's) we illustrated in detail in the beginning of this chapter (p. 73). Although that is a chemical, and not a merely physical fact, the molecular hypothesis can explain it also. Since matter is composed of molecules, a chemical action between two kinds of matter must consist really in an interaction of the molecules of each kind. Molecules of the two kinds must meet and they may then either combine to form a compound molecule, or they may exchange material in some fashion. Since equal volumes are often the exact quantities required for the action, it appears most likely that in equal volumes of different gases (at the same temperature and pressure) the numbers of mole- cules present are equal. This addition to the molecular hypothesis was first suggested by an Italian physicist, Avogadro (1811). When two gases interact in equal volumes (like hydrogen and chlorine), one molecule of each is all that is required for a small sample of the change in question. Since two volumes of 84 TEXT BOOK OF CHEMISTRY hydrogen are required to unite with one volume of oxygen (p. 78), clearly the interaction involves two molecules of hydrogen for every one of oxygen. Since two volumes of steam are produced, evidently the two molecules of hydrogen and one of oxygen yield two molecules of water. This hypothesis of Avogadro helps us to a clearer notion of how these chemical changes take place. The idea that all gases contain equal numbers of molecules in equal volumes also explains why all gases behave alike when subjected to equal pressures, or to equal changes in temperature (Boyle's and Charles' laws, p. 44). No facts which conflict with this hypothesis are known, and all the known facts confirm it. "Hence, Avogadro's hypothesis has been accepted by chemists, and since 1858 has been the keystone of chemical theory. Consequences of Avogadro's Hypothesis Molecular Weights. Equal volumes of the same gas (at the same tem- perature and pressure) have equal weights. But equal volumes of different gases have different weights. The differences are often very great. Thus, bulk for bulk, oxygen is sixteen times as heavy as hydrogen, and hydrogen iodide sixty-four times as heavy. Now, if equal volumes of different gases' contain equal numbers of molecules, these differences must be due to the differing weights of the several kinds of molecules. Thus, measuring the weights of equal volumes of different gases will give us the relative weights of their molecules. For example, since 22.4 1. of oxygen weigh 32 g. (p. 75), while the same volume of hydrogen weighs 2.016 g. and of hydrogen chloride 36.468 g., and these are the weights of equal numbers of molecules, the individual molecules must differ in weight in the ratio 32 : 2.016 : 36.468. These are the relative weights of the three kinds of molecules. In chemistry the weights of 22.4 1. (at and 760 mm.) of various gases are called the molec- ular weights of those gases. The unit quantities of various sub- stances are therefore spoken of, technically, as the molecular CHEMICAL UNITS OF WEIGHT. FORMULA 85 weights of those substances. The unit volume, 22.4 1., is called the gram-molecular volume (G.M.V.). The number of molecules actually contained in the G.M.V. is approximately 6 followed by 23 zeros, which we may write 6623. The gram-molecular weight of any gas is therefore the weight in grams of 6623 molecules of the gas. Molecular Weight, Measurement of. The molecular weight is measured by weighing any convenient volume of the gas (say 200 c.c.), and calculating by proportion the weight of 22.41. If the gas or vapor was not measured at and 760 mm., the measured volume must be reduced by rule to standard conditions before the weight of 22.4 1. is calculated.* For example, 190 c.c. of a gas at and 760 mm. weigh 1.23 g. If x be the weight of 22.4 liters (= 22,400 c.c.), 190 : 1.23 :: 22,400 : x (= 145 g.). Again, 210 c.c. of a vapor at 100 and 743 mm. weigh 1.12 g. This volume at and 760 mm. would become: and 150.3 : 1.12 :: 22,400 :x(= 167 g.). Molecular Weight and Density. Density is the term used in physics for the weight of 1 cubic centimeter of a substance. Thus, the density of water at 4 C. is 1, because at 4 C., 1 c.c. of water weighs 1 g. The density of ammonia gas (wt. of 1 c.c. at and 760 mm.) is 0.000759 g. Since the molecular weight is the weight of 22.4 1., or 22,400 c.c., the molecular weight of a gas is obtained by multiplying the * In practice, owing to the fact that Boyle's and Charles' laws do not describe the behavior of any known gas exactly (they apply only to a "perfect" gas), certain additional, small corrections have to be applied when very precise values are required. 86 TEXT BOOK OF CHEMISTRY density (if known) by 22,400. Thus, the molecular weight of ammonia is 0.000759 X 22,400 = 17.0. Densities are often given on the scale, density-of-air = 1. Now 22.4 liters of air weigh 28.95 g. If a gas has a density twice that of air, 22.4 liters of this gas would weigh 28.95 X 2 g. For example, the density of carbon dioxide (air = 1) is 1.52. The molecular weight is therefore 28.95 X 1.52, or 44.0.* Atomic Weights. Since a compound substance can be formed by union of the elementary substances, and decomposed to give these substances, its molecule may be assumed to contain those constituent elements as distinct parts of the mass. Those elementary parts of a molecule are called atoms. When two ele- mentary substances combine, the process involves the union of the two kinds of atoms to form compound molecules. Now, in the table (p. 77), we recorded the weights of the constituent ele- ments which together made up the molecular weights of the com- pounds. Then, upon examining the weights of any one element contained in molecular weights of different compounds, we saw (p. 77) that one (the smallest) could be taken as the unit weight, of which the others were multiples by some whole number. Ob- viously then, if the molecular weights are the relative weights of different kinds of molecules, the unit weights of the elements (35.46 for chlorine, etc.) are the relative weights of the atoms of the different elements. The relative weights of the different kinds of atoms, such as 1.008 for hydrogen, 16 for oxygen, 35.46 for chlorine, are called the atomic weights of the respective ele- ments. These weights, as we have seen (p. 80), are relative to the weight of the atom of oxygen, when the weight of the latter is taken as 16, and the weight of the molecule of oxygen is taken as 32, A compound molecule may contain one, or more than one, atom of each of the elements forming the compound. In the molecular * This weight of carbon dioxide contains 12 parts of carbon (one atomic weight) and 32 parts of oxygen (two atomic weights) total 44. CHEMICAL UNITS OF WEIGHT. FORMULAE 87 weight of compounds, weights of elements which are smaller than the atomic weight (p. 77), are not found to occur (see e.g., p. 86, footnote). This indicates that, in chemical change, fractions of atoms play no part. The name atom (Greek, not cut, or not divided) records this fact. The Atomic Hypothesis and Definite Proportions. The idea of atoms furnishes at once an explanation of the law of definite proportions. Evidently, every molecule of a given substance must always contain the same numbers and the same kinds of atoms, so that the proportions by weight of the constituents of the compound as a whole must be alike in all samples. Dulong and Petit' 's Law. When an element gives no volatile compounds, as is the case with calcium, we can still find the atomic weight. We first find for the element an equivalent weight (p. 53). Thus, the weight of calcium which combines with 35.46 parts of chlorine is 20 parts. This weight, or some multiple of it by a whole number, does in fact always correctly represent the amount of calcium combined with atomic weights of elements whose atomic weights are known. Hence 20, or some multiple of 20 is the required atomic weight. Dulong and Petit (1819) dis- covered a rule by which the correctness of such an atomic weight could be checked. When the atomic weight of an element is multi- plied by the specific heat of the element in the solid form, the prod- uct is roughly 6.4. Thus, the known atomic weight of magnes- ium is 24.3, its specific heat 0.245, and the product is 5.95. Again, the known atomic weight of mercury is 200.6, and its specific heat 0.0335, and the product 6.7. Now the specific heat of calcium is 0.170, and the product 20 X 0.170 is equal to 3.4. But 2 X 20 X 0.170 = 6.8. The value 2 X 20, or 40 for the atomic weight is therefore the correct one. The following table contains additional illustrations of this law: 88 TEXT BOOK OF CHEMISTRY Element. Atomic Wt. Sp. Ht. Prod- uct. Element. Atomic Wt. Sp. Ht. Prod- uct. Lithium 7 94 6.6 Iron 56 0.112 6.3 Sodium 23 29 6.7 Zinc 65.4 0.093 6.1 Magnesium. . Silicon . . . Phosphorus (Yellow) . . Calcium . . . 24.3 28.3 31 40 0.245 0.16 0.19 0.170 6.0 4.5 5.9 6.8 Bromine (Solid) Gold Mercury (Solid) Uranium . . . 80 197 200 238.5 0.084 0.032 0.0335 0.0276 6.7 6.3 6.7 6.6 Another way of expressing this law will give it greater chemical significance. The specific heats are the amounts of heat required to raise one gram, that is one physical unit, of each element through one degree. When we multiply this by the atomic weight, we obtain the amount of heat required to raise one gram-atomic weight of the element, that is, one chemical unit, through one degree. The values of this product are approximately equal. Since there are equal numbers of atoms in one gram-atomic weight of each ele- ment, it follows that : Equal amounts of heat raise equal number of atoms of all elements in the solid form through equal intervals of temperature. RELATIONS BETWEEN THE STRUCTURE AND BEHAVIOR OF MATTER We have seen that matter is composed of minute particles called molecules. Just as we can thoroughly understand the be- havior of a watch or an automobile engine only if we know the details of its structure, and how the parts work, so we can under- stand the physical and chemical behavior of matter in masses only if we are familiar with its ultimate mechanism. Hence, we must now take up the structure of matter in its three states, the gaseous, the liquid, and the solid. In doing this, we shall keep constantly in view the connection between the molecular relations and the general behavior of the matter. RELATIONS BETWEEN STRUCTURE AND BEHAVIOR 89 The Molecular Structure of Gases. The most noticeable fact about gases is that they can be compressed to an enormous extent. Oxygen at 760 mm., for example, can be reduced by pressure to one two-hundredth of its volume, or even less. The compression does not affect the individual molecules, and there- fore does not diminish the volume actually occupied by the oxygen, but it crowds the molecules closer together and diminishes to one two-hundredth the space between them. Compressing a gas is, in fact, mainly compressing the empty space of which it chiefly consists. To understand what follows, the reader must keep constantly before him a mental image of a jar of gas as consisting of small particles separated by relatively wide, empty spaces. The molecules are in rapid motion and move in straight lines, excepting when they strike one another or the walls of the The Properties of Gases. Let us now note the more obvious qualities of gases, printing in italics the fact concerning a mass of gas and in black type the property of the molecules which accounts for the fact. The most remarkable thing about a gas, considering the loose- ness with which its material is packed, is the total absence in it of any tendency to settling or subsidence. Since the molecules cannot be at rest upon one another, as the great compressibility shows, we are driven to conclude that they are widely separated from one another, and that they occupy the space, otherwise a complete vacuum, by constantly moving about in all directions. But a moving aggregate of particles which does not even finally settle must be in perpetual motion. We must, therefore, believe the molecules to be wholly unlike particles of matter in having perfect elasticity, in consequence of which they undergo no loss of energy after a collision. They must continually strike the walls of the vessel and one another and rebound, yet without loss of motion. The fact that each gas is homogeneous, efforts to sift out lighter or 90 TEXT BOOK OF CHEMISTRY heavier samples having failed, requires the supposition that all the molecules of a pure gas are closely alike. The diffusibility of gases is due to the motion of the molecules, and their permeability to the space available to receive molecules of another gas. These two modes of behavior involve no additional molecular properties. The word " diffusion" is often thought to mean the property of a given mass of gas in virtue of which another gas can mix with the given mass. This property is not diffusibility but permeability. It is the other gas, which makes its way into the given gas, which is diffusing. Diffusion is spontaneous motion of the parts of a gas away from their original location. Unless this motion is into an empty space, the diffusing molecules must, of course, move into another body of gas. In the case of the jars of hydrogen and air, each gas moved in part out of its original jar (diffused), and each received parts of the other gas into its jar (was permeated). Boyle 9 s Law and Charles 9 Law. Passing now to Boyle's law (p. 44), the thing to be accounted for is that when a sample of a gas diminishes in volume, its pressure increases in the same -proportion. Let the diagram (Fig. 37) repre- sent a cylinder with a movable piston, upon which weights may be placed to resist the pressure. Now Tthe pressure exercised by the gas under the piston cannot be like the pressure of the hand upon a table, since we have just assumed that the particles are not even approximately at rest, and the spaces between FIG 37 them are enormous compared with the size of the mole- cules themselves. The gaseous pressure must there- fore be attributed to the colossal hailstorm which their innumerable impacts upon the piston produce. If this is the case, the com- pressing of a gas must consist simply in moving the partition downwards, so that the particles as they fly about are gradually restricted to a smaller and smaller space. Their paths become on RELATIONS BETWEEN STRUCTURE AND BEHAVIOR 91 an average shorter and shorter. Their impacts upon the walls be- come more and more frequent. So the pressure which this causes becomes greater and greater, and is proportional to the degree of crowding (the concentration) of the molecules. There are two other points to be added. When we diminish the volume to one-half, we find from experience that the pressure be- comes exactly, or almost exactly, twice as great. This must mean that, although the particles are becoming crowded, they do not interfere with one another's motion, excepting of course where actual collision causes a rebound. Only in the absence of inter- ference would doubling the number of molecules per unit of volume give exactly double the number of impacts on the walls. Hence the molecules must have practically no tendency to cohesion. Finally, the molecules must be supposed to move in straight lines between collisions* A baseless idea, that the molecules of a gas repel one another, still lingers in some quarters. There is no evidence of this. The molecules pay almost no attention to one another, either by attraction or repulsion. Boyle's law therefore adds four more details concerning molec- ular behavior, namely, that the impacts of the particles produce the pressure, that the crowding of the molecules represents the con- centration of the material and that the particles move in straight lines and show almost no cohesion, since pressure and concentration are very closely proportional to one another. How, now, can we account for Charles J law (p. 45), according to which an increase in pressure (or in volume) results from heating a mass of rapidly moving molecules? The action of a particle colliding with a surface is measured in physics in terms of its mass and its velocity. It is evident that heating a cloud of mole- cules would not increase the mass of each, and it must therefore increase the velocity of each, since the kinetic energy of all becomes greater. 92 TEXT BOOK OF CHEMISTRY Avogadro's Law. The identical general behavior of all kinds of gases suggests that their structures may be all alike. Avogadro (1811), the professor of physics in Turin, put forward the hypothe- sis that the numbers of molecules in equal volumes of different gases, at the same temperature and pressure, might be equal. A more strict study of the properties we have been considering, and of some additional facts, has since shown that no other con- jecture than Avogadro's would be consistent with them. Thus it is now accepted as a fact, and is known as Avogadro's law. It may also be put in the form: At the same temperature and pressure, the molecular concentration of all kinds of gases has the same value. Diffusion. The law of diffusion (p. 90) harmonizes with the conceptions of molecular structure without further additions to the latter. The speed of the hydrogen molecule at room temperature is 1840 meters per second. The masses of the hydrogen and oxygen molecules are as 1 : 16, and the speeds of diffusion are as Vl6 : Vl, or 4 : 1. Hence the speed of the oxygen molecule is one-fourth of 1840, or 460 m. per sec. Calculation shows the activity of the molecules to be such that, in air, the number striking a single square centimeter of surface per second would fill no less than twenty liters. Liquefaction of Gases: Critical Temperature. Finally, gases can be liquefied by sufficient cooling and compression. This fact compels us to suppose that, after all, even gaseous molecules have a tendency to cohesion. This cohesion is scarcely perceptible so long as the gas is warm and is diffuse. Thus, 2 liters of oxygen at 1 atmosphere pressure, when subjected to 2 atmospheres pressure, give 0.9991 liters instead of 1 liter. The additional contraction of 0.0009 liters (0.9 c.c.) is due to the effect of cohesion when the molecules are thus crowded closer together. The gases which are more easily liquefied than is oxygen show greater effects. RELATIONS BETWEEN STRUCTURE AND BEHAVIOR 93 Thus, 2 liters of sulphur dioxide at 760 mm., when subjected to 2 atmospheres pressure, give only 0.974 liters, showing a contraction due to cohesion of 26 c.c. These data refer to 0. At lower tem- peratures the contractions due to cohesion become rapidly greater. This cohesion is not of the nature of gravitational attraction. We can readily understand, therefore, that when the kinetic energy of the molecules is sufficiently reduced by cooling (namely, to, or below the critical temperature, see below), and the molecules are brought sufficiently close together, the tendency of the mole- cules to cohere causes the gas to condense and assume the liquid form. In 1869 Andrews found that carbon dioxide could be liquefied at by 38 atmospheres pressure, and at 30 by 71 atmos- pheres, but that above 31.35 it could not be liquefied by any pressure. He discovered that each gas has a critical >*. temperature, as he called it. For carbon dioxide, this temperature can be observed by placing a heavy-walled, glass tube (Fig. 38), half -filled with liquid carbon di- oxide, in a beaker of water, and gradually raising the temperature of the latter. At 31.35, the surface be- tween the liquid and gas becomes hazy and vanishes. When the temperature falls once more, the surface re- appears at 31.35. This shows that, with Faraday's "permanent" gases, a temperature below the critical FlG ~ 38 point had not been employed. The critical temperature of oxygen is 118, of hydrogen 234, of carbon dioxide 31.35, of sulphur dioxide 156, of water 358. Another Deviation from the Laws of Gases. A Perfect Gas. It may be added that when a gas is already under very high pressure, and very closely packed, an increase in the pressure does not produce quite as great a diminution in volume as Boyle's law leads us to expect. This reminds us that we are only the space between the molecules, and not the volumes of the molecules themselves, and therefore not the total volume of the 94 TEXT BOOK OF CHEMISTRY gas. When, on severe compression, the volume occupied by the molecules themselves has become an appreciable fraction of the whole volume, additional compression does not affect the whole volume, and the contraction is smaller than Boyle's law would indicate. Thus, 2 liters of hydrogen, even at one atmosphere pressure, when subjected to two atmospheres pressure, give 1.0006 liters, instead of 1 liter. The last two effects (namely, those due to the tendency to cohesion of, and the space occupied by the molecules) are called deviations from the laws of gases. In consequence of these in- dividual deviations, there are not exactly equal numbers of mole- cules in equal volumes of any two different gases, at the same temperature and pressure. An imaginary gas, which exhibits neither deviation, called a perfect gas, is often referred to in discussing the behavior of gases. Summary. We may now summarize the principal facts about gases in mass, appearing in italics above, with the corre- sponding features of the molecular relations, in heavy type, which we have added one by one. Facts About Gases in Mass. Corresponding Relations of Molecules. Compressibility Diffusibility . . Permeability Non-settling . . Homogeneity . Pressure . . . Boyle's law . . Charles' law Above and other facts Law of diffusion . . . Gases can be liquefied Vacuum + molecules widely separated. Molecules in rapid motion. Empty space relatively large. Molecules perfectly elastic. Molecules of any one substance closely alike Due to impacts of molecules. Pressure proportional to concentration of the molecules. Molecules move in straight lines and, when widely scattered, show no tendency to cohesion or to repulsion. Rise in temperature increases the velocity, and therefore the kinetic energy of the molecules. Avogadro's law. Molecules do possess a tendency to cohesion, which becomes conspicuous when they are cooled and closely crowded together. RELATIONS BETWEEN STRUCTURE AND BEHAVIOR 95 History of the Kinetic Molecular Theory. This theory was first suggested by Daniel Bernoulli (1738), who explained by its means the pressure, and compressibility of gases. Lomonossov (1748) developed the theory very completely and by means of it, explained Boyle's law and the effects of changes in temperature. He also anticipated from the theory the existence of the second deviation from the law of gases (1749), a discovery usually cred- ited to Dupre (1864). He likewise pointed out that there was no limit to the maximum velocity of a molecule, and therefore no upper limit of temperature, but that there must be a lower limit (the absolute zero) at which the molecules would be at rest (1744). This work was entirely forgotten, until attention was called to it in 1904 by Menschutkin. Similar views were expressed by Waterston (1845), but were still so much ahead of the time that the committee of the Royal Society did not approve the paper for publication, and it was dis- covered in the archives of the society, long afterwards, by Lord Rayleigh. The development of the theory, so far as it applies to heat is therefore credited to Joule (1855-60) and, in respect to all properties of gases, to Kronig (1856) and Clausius (1857), who knew nothing of the earlier work. Molecular Relations in Liquids. The fact that even great pressures produce little diminution hi the volume of a liquid shows that the free space, present in gases, is absent in liquids. The measured effect^ of various pressures show, for example in the case of water, that to reduce the volume to one-half would require, not doubling the pressure as in a gas, but increasing it from 1 to 10,000 atmospheres. The molecules of a liquid are actually in contact with one another. The phenomena connected with surface tension, such as co- herence into drops, show that cohesion plays a much larger part in liquids than in gases. On the other hand, liquids which are capable of mixing (e.g., alcohol and water), when placed above one 96 TEXT BOOK OF CHEMISTRY another in the same vessel, do mix, slowly, by diffusion. This indicates that motion of the molecules, although much impeded by friction, has not been annihilated by cohesion. The escape of vapor that is, of part of the liquid in gaseous form likewise proves that the molecules in the liquid are in motion. The rela- tions of liquid and vapor can be discussed most effectively in the next chapter, in connection with the case of water and steam. Molecular Relations in Solids. The properties of solids differ from those of liquids chiefly in the fact that the solid has a definite form of which it can be deprived only with difficulty. This we may explain in accordance with the kinetic hypothesis by the supposition that the cohesion in solids is very much more prominent than in liquids. We obtain solids from liquids by cooling them; in other words, by diminishing the kinetic energy and therefore the velocity of the particles. The cohesive tendency of the latter is thus able to make itself felt to a greater extent. If, conversely, we heat a solid, or, according to the hypothesis, if we increase the speed with which the particles move, the body first melts and gives a liquid, and this finally boils and becomes a gas. The intrinsic cohesion of the particular substance can undergo no change, but the increasing kinetic energy of the particles steadily and continuously obliterates its effects. Yet some motion still survives in a solid. Thus we find that when the layer of silver is stripped from a very old piece of electroplate, the presence of this metal in the German silver or copper basis of the article is easily demonstrated. The tendency of all solids to assume crystalline forms, which show definite cleavage and other evidences of structure, distin- guishes them sharply from liquids. The force of cohesion in liquids is exercised equally in different directions. In solids it must differ in different directions in order that structure may re- sult. Since each substance shows an individual structure of its own, these directive forces must have special values in magnitude and direction in each substance. RELATIONS BETWEEN STRUCTURE AND BEHAVIOR 97 Exercises. 1. State the change which takes place in the total volume of the gases or vapors in each of the six actions mentioned, p. 73. 2. If 1 liter of oxygen at and 760 mm. weighs 1.429 g., what is the molecular weight (pp. 75, 85) ? 3. Calculate the molecular weight of a gas, 200 c.c. of which at and 760 mm. weigh 2.1 g. (p. 85). . 4. Find the molecular weight of a gas, of which 250 c.c. at 18 and 752 mm. weigh 2.5 g. (p. 85). 5. The following are the weights of sulphur contained in the molecular weights of several of its compounds: 32.06, 64.12, 96.18. What is the atomic weight of sulphur (pp. 75, 77) ? 6. What information is contained in each of the formulae: CS2, PCls, A1 2 O 3 (p. 86)? 7. If two gases combine in the ratio 5 : 2 by volume, in what relative numbers do their molecules interact (p. 84) ? 8. The density of sulphur dioxide is 0.000286. What is its molecular weight (p. 85) ? 9. The density of a gas, air = 1, is 2.3. What is its molecular weight (p. 86) ? 10. An element combines with hydrogen in the proportion 10.03 : 1, and its specific heat is 0.2. What is its atomic weight (p. 87)? 11. The molecular formula of a gas is CH 4 . What is the gram- molecular weight? What volume does this weight occupy? What is the weight of 1 liter of the gas? CHAPTER VIII MAKING OF FORMUUE AND EQUATIONS THE formula (p. 81) is a condensed statement of the composition of a substance. Before we can make (i.e., calculate) the formula for a substance, we must (1) measure the proportions by weight of the constituent elements. Then, we must (2) express these propor- tions in multiples of the known atomic weights of the elements. Analysis and Synthesis. In the case of water H^O we saw (p. 59) how the weights of hydrogen and of oxygen required to give a measured amount of water were determined. The com- position of the water was found out by putting the substance to- gether out of the elements. This method is called synthesis (Greek, putting together). In the case of mercuric oxide we can take a weighed amount of the oxide, decompose it, and weigh the mercury formed. The difference is the weight of the oxygen. This process, of decom- posing a substance to learn its composition, is called analysis, the Greek word for decomposition. One or other sometimes both of these plans can be used with every compound. Some of the results of such experiments have been given in the earlier chapters. For example: Tin (100) + Oxygen (26.89) -> Tin oxide. Lead (100) + Oxygen (7.72) -> Lead oxide (p. 10). Iron (100) -f Oxygen (43) -> Ferric oxide (p. 10). Zinc (2.04) + Sulphur (1) -> Zinc sulphide. Mercuric oxide (108) -> Mercury (100) + Oxygen (8) (pp. 15, 19). MAKING OF FORMULA AND EQUATIONS 99 Making Formulae. In the formulae, these proportions are to be replaced by multiples of the atomic weights by whole num- bers. We therefore divide the quantity of each element by the corresponding atomic weight. This operation gives us the factors by which the atomic weights are to be multiplied. The atomic weights we find in the table, where the values determined by the most expert chemists are collected. For example. In the case of tin oxide the proportion of tin to 100 oxygen is . The atomic weights are 119 and 16, respectively. Ao.oy 1 rin 100 -r 119 = 0.84, and 26.89 -5- 16 = 1.68. The proportion- 1 ^- 26.89 , 119 X 0.84 now becom is - 16 X 1.68 Now this proportion like all chemical proportions must be expressed in multiples of the atomic weight by whole numbers. Hence, we next find the greatest common measure of the two factors. It is 0.84. Indeed, in this simple instance, we can see that the ratio of the factors is 1 : 2. Dividing above and below ^ 119 X 1 by 0.84, we get 16 x 2 Now, the symbols stand for the atomic weights. Substi- tuting these symbols, the proportion becomes - . The \J X formula is therefore SnC>2. Applying the same process to the lead oxide, we get 100 207.1 X 0.483 207.1 X 1 Pb X 1 7.72 16 X 0.483 16 X 1 Oxl , or PbO. Treating the other data in the same manner, we find : ferric oxide F^Os, zinc sulphide ZnS and mercuric oxide HgO. If the composition of the substance is given in percentages, the same process is used. Thus, the case of sodium sulphate works out as follows: 100 TEXT BOOK OF CHEMISTRY Elements. Percentages. At. wt. Quotient 4- Formula. Sodium 32.43 23 X 1.41 0.705 Na X 2 Sulphur 22.55 32 X 0.705 0.705 S Oxygen 45.02 16 X 2.814 0.705 O X 4 The formula is therefore Making Equations. The condensed statements of chemical changes which we have been using can now be still further sim- plified by using the formula in place of the names of the sub- stances (p. 81). Thus Sn + O 2 -SnO 2 This is to be read: 119 parts (or 1 atomic weight) cf tin, acting chemically with 2 X 16 parts (or two atomic weights) of oxygen, give 157 parts of stannic oxide. We may also read it thus: 1 atom of tin with 1 molecule of oxygen give 1 molecule of stannic oxide. In making an equation there are four stages or steps : 1. Find out by experiment what the substances used and pro- duced are. 2. Learn the molecular formula of each substance. 3. Set down the molecular formulae in the form of a skeleton equation. Place the formulae of the initial substances on the left, and those of the products on the right. 4. Adjust, or balance the equation. For example: (1) When hydrogen and oxygen combine, water is formed. (2) The molecular formulae are H 2 , O 2 , and H 2 O. (3) Skeleton equation: H 2 + 2 -> H 2 O. (4) In accordance with the law of conservation of mass, the numbers of atomic weights (or atoms) of each element must be the same after the action as before it. Now the skeleton equation shows two atomic weights of oxygen before, and, thus far only one after the action, whereas there ought to be two there also. With Oa (2 X 16 parts) we have enough oxygen to give 2H 2 0, which MAKING OF FORMULA AND EQUATIONS 101 contains 2 X 16 parts of oxygen. But this will require us to take 2H 2 to " balance" the equation. The final equation is, therefore: Balanced Equation: 2H 2 -f 2 -> 2H 2 0. Observe, we could not adjust the difficulty by writing H 2 + O H 2 0, because each substance must be represented by its molec- ular formula, which stands for the weight of the substance in the standard volume of 22.4 liters, or one chemical unit weight, and in the case of oxygen this is O- (= 32 g.). Putting this in terms of the hypothesis, each formula must represent 1 molecule, and the molecules of oxygen contain 2 atoms. Hence we could not divide the oxygen molecule. But we could take more than one molecule of hydrogen, so we took 2 molecules of this substance. The coefficients in front of formulae multiply the whole formula. 2H 2 is equivalent to 2(H 2 O), or two whole molecules of water. Balancing Equations. Learning to balance equations correctly comes only by practice. Take, again, the case of iron rusting. The substances are iron (Fe), oxygen (02) and ferric oxide (FeaOs). The skeleton equation is We are not permitted to alter these formulae themselves, but we may put coefficients in front of any of them to make the number of atomic weights alike on both sides. A good rule is to pick out the largest formula and reason back from that. Here, this is Fe-jOs. To get the Os in threes, we must clearly take 3O 2 ( = 6O). That will give us 2F 2 O 3 . This, in turn, will require 4Fe: Balanced: 4Fe + 3O 2 - 2Fe 2 Oa. Equations for Actions Already Studied. In the prepa- ration of oxygen (p. 15) we used mercuric oxide and got mercury and oxygen: Skeleton: HgO-> Hg -f 2 . Balanced: 2HgO -> 2Hg + O 2 . 102 TEXT BOOK' OF CHEMISTRY Potassium chlorate has a composition shown by the formula KC1O 3 . It gives (p. 30) potassium chloride (KC1) and oxygen (0 2 ). Skeleton: KC10 3 -> KC1 + O 2 . Balanced: 2KC10 3 - 2KC1 + 3O 2 . The variety of chemical change, where one substance gives two (or more) substances, decomposition (p. 16), is readily recognized in these equations. Substances, like the manganese dioxide (catalytic agent) used here, and the water so often employed as a solvent, when they undergo no chemical change, are omitted from the equation. They are simply part of the apparatus. When the water takes part in the action, however, it must, of course, be included. Thus sodium peroxide (Na 2 O 2 ) and water (H 2 O) interact (p. 31) to give sodium hydroxide and oxygen: 2Na20 2 + 2H 2 O -> 4NaOH + 2 . The Preparation of Hydrogen (p. 50) from sodium (Na) and water gives sodium hydroxide (NaOH) and hydrogen (H 2 ) : 2Na + 2H 2 -> 2NaOH + H 2 . When steam is passed over iron (p. 51), we get hydrogen and magnetic oxide of iron (Fe 3 4 ) : The liberation of hydrogen by the action of zinc (Zn) upon sulphuric acid (H 2 SO 4 ), where the products (p. 52) are hydrogen and zinc sulphate, is shown thus: Zn + H 2 SO 4 -* H 2 + ZnSO 4 . Again, iron and hydrochloric acid (HC1) give hydrogen and ferrous chloride (FeCl ? ): Fe + 2HCi-H 2 -t-FeCl 2 , MAKING OF FORMULA AND EQUATIONS 103 [n the last two equations the variety of chemical changes called displacement (p. 51), where one elementary substance displaces another from a compound, is well illustrated. The equation for the formation of water by union of hydrogen and oxygen, 2H 2 + O 2 ->2H 2 O, has already (p. 60) been discussed. The reduction of an oxide, such as magnetic oxide of iron or cupric oxide, by hydrogen (p. 57), gives the metal and water: 4H 2 < 3Fe + 4H 2 O. Upon examining these equations for reductions, we perceive that they are illustrations of displacement also. Reversible Actions (p. 68), like the decomposition of water by heating, and the recombination of the elements on cooling (p. 68), are shown by using two arrows: The equation may be read from either end. The decomposition and formation of hydrates (p. 69) are also reversible actions. In the case of zinc sulphate, the equation is ZnSO 4 -f 7H 2 O<=ZnSO 4 , 7H 2 O. Reaction Formulae. In the foregoing formula for the hydrate of zinc sulphate, it will be seen that we do not add together all the atoms of oxygen, and write ZnHuSOn. The latter would show the composition of the substance correctly, but it would show nothing more. Now chemists find it convenient, frequently, to alter the formula so that it shall indicate also some important chemical property or reaction of the substance. Hence the formula ZnS0 4 , 7H 2 0, which indicates at a glance the relationship of the substance to zinc sulphate (ZnSO 4 ). The hydrate is made from zinc sulphate by adding water, and is easily decomposed into these 104 TEXT BOOK OF CHEMISTRY two substances again. The reaction formula hints at this familiar reaction. Note, however, that the comma ( , ) does not indicate a mixture of the materials, such as ZnSCX and H 2 0, but a single substance composed of both. The plus ( + ) sign is used between the formulae of different, uncombined substances in a mixture. In accordance with this plan, washing soda, hydrate of sodium carbonate (p. 69), is written Na2C0 3 ,10H 2 O, and bluestone (p. 69), hydrate of cupric sulphate, CuSC>4, 5H 2 O. Dissociation. A decomposition which, like that in p. 68, is reversible is called a dissociation. When heat is the agent pro- ducing the change, it is sometimes called a thermal dissociation. Not all decompositions are reversible. Thus potassium chlorate decomposes to give potassium chloride and oxygen, but these products will not combine under any known conditions, directly, to give potassium chlorate. Molecular Formulse. In this chapter, for the sake of simplicity, we have so far left in the background the fact that the formula must represent a molecular weight of the substance, as well as its composition. The total weight, for which the symbols in a formula stand, must be equal to the weight of the substance occupying the gram-molecular volume. In other words, the formula-weight must represent one cube-full (Fig. 36, p. 75) of the substance. This is true, as we have seen, of the formulas H 2 and HC1 (see p. 75). In the cases of tin oxide (Sn0 2 ) and ferric oxide (Fe 2 O 3 ), we have substances which cannot be converted into vapor or dissolved, so that their molecular weights are unknown. In such cases, as we have seen (p. 78), we use the simplest formula that will show the correct proportions. Molecular Formulae of Simple Substances. With oxy- gen (O 2 ) and hydrogen (H 2 ), however, the double formulae are used, instead of the simpler and H, because the weights 2 X 16 MAKING OF FORMULAE AND EQUATIONS 105 and 2 X 1.008 are the ones which fill the cube. The molecules of all elements are not diatomic, however. Thus the cube-full of mercury vapor weighs only 200.6, the same as the atomic weight, and the correct molecular formula of the element is therefore Hg. Similarly, the correct formulae are Na (sodium), K (potassium), and Zn (zinc). But the weight of 22.4 liters requires us to write N 2 for nitrogen, P 4 for phosphorus vapor and 82 for sulphur vapor (above 1000). Hence the equation reads: 2Zn + 82 -> 2ZnS. After all, there is nothing surprising in the fact that the mole- cules even of elementary substances should, in some cases, contain several atoms. All that it means, in the case of an element with diatomic molecules, such as oxygen (0 2 = 32), is that, when oxygen combines with another element, each molecule of oxygen will be divided between two molecules of the product if the latter contain only 16 parts of oxygen each. Molecular Formulae of Compounds. The need of atten- tion to making our formulae molecular comes out also in the cases of many compounds. Thus, formaldehyde (a disinfectant) has the composition CH 2 O, and its molecular weight is 30, so that CH 2 O (= 12 + 2+16) is the correct formula. But acetic acid (the sour substance in vinegar) has the same composition, CH 2 O, only its molecular weight is 60, and the formula is therefore written C 2 H 4 O 3 (= 24 + 4 + 32). For gaseous and volatile substances the correct molecular for- mulae are always used. Thus, the equation is written: P 4 Oi , because the molecular weight of phosphorus pentoxide is 284, and requires the formula P4Oio rather than P 2 C>5. Warnings. Do not place the formulae of the products upon the left, but upon the right-hand side of the equation. Point the arrow towards the products. Use the molecular formulae for elementary substances (O 2 , H 2 , N 2 , etc.). The molecular formulae, when they are known, are the 106 TEXT BOOK OF CHEMISTRY only ones given in the text. The symbols, as given in the table on the rear cover, must be made into molecular formulae before use in equations. Exercises. 1. Using the data given on p. 73, calculate the formulae of sulphur dioxide, phosphorus pentoxide, and carbon dioxide. 2. Using the results of 1, make and balance the equations for the union of each of the three elements, sulphur, phosphorus, and carbon, with oxygen. 3. Using the information on p. 36, make and balance the equations for the interaction with oxygen of: (a) carbon disul- phide (082); (b) zinc sulphide (zinc oxide is ZnO), and (c) wood (assuming the formula of the latter to be that of cellulose, C 6 Hi 5 ). 4. Make and balance the equations for the actions on water of: (a) potassium (giving KOH), and (6) calcium (giving CaO 2 H 2 , usually written Ca(OH) 2 ). 5. Using the data in regard to the action of zinc on hydro- chloric acid, given on p. 53, calculate the formula of zinc chloride and make the equation. 6. Make equations for the action of magnesium and aluminium upon hydrochloric acid (giving MgCl 2 and A1C1 3 ) and upon sul- phuric acid (giving MgSO 4 and A1 2 (SO 4 ) 3 ). 7. Make a molecular equation for the decomposition of hydro- chloric acid by electrolysis (p. 55). 8. Make equations for the formation of the hydrates of sodium carbonate (NaaCOa, 10H 2 O) and of cupric sulphate (CuS0 4 , 5H 2 0), by union of the anhydrous substances with water. CHAPTER IX SOLUTION THE property that many substances have of dissolving hi others is a most interesting and valuable one. The value lies chiefly in the fact that some substances are easily soluble in a given liquid and others are, practically, not soluble in it at all. These differ- ences in solubility enable us to accomplish many things otherwise impossible. Thus, we separated sulphur from iron (p. 13), by using carbon disulphide (082) to dissolve the former. We must first learn precisely what is meant by a solution, and then we shall be ready to understand the uses and properties of solvents and solutions. Solution. We distinguish carefully between a solution and a mere mixture. In a solution the dissolved substance is so completely spread throughout the solvent (such as water) that it gives no evidence to the eye of its presence. The only exception to this is when the substance is colored, and confers a general tint on the whole solu- tion. Furthermore, the dissolved substance never settles out, no matter how long we wait. On the other hand, when the liquid is dull, or cloudy, or muddy, the liquid is a mixture. The substance is simply suspended, not dissolved, and will eventually settle out. Sand, shaken with water, settles at once. Flour, mixed with water, settles more slowly. Filtration (p. 13) quickly separates the flour, which remains on the paper, from the water, which runs through. These are not dissolved, but suspended substances. Milk owes its white appearance to minute drops of oily matter which reflect much light from their surfaces. They pass easily 107 108 TEXT BOOK OF CHEMISTRY through filter paper. But they slowly rise to the top, being lighter than water, and are likewise not dissolved, but suspended. A mixture of two liquids, like milk, is called an emulsion. Chemists commonly call the dissolved body the solute. Gases, liquids, and solids may all be solutes, and dissolve in suitable solvents. Solvents. Water is by far the commonest and most useful solvent. Very many inorganic substances dissolve in it easily. The fact that many (like zinc and sand) do not, enable us to separate the components of a mixture containing a soluble and an insoluble substance. Many organic substances, such as fats, paraffin, petroleum, tar, rubber, cotton, paper, shellac, and so forth, do not dissolve to any measurable extent in water. But fats dissolve readily in ether (C^ioO), in carbon disulphide (82), in carbon tetrachloride (CC1 4 ), and in chloroform (CHC1 3 ). For this reason these sub- stances remove grease which has accidentally got into cloth. Paraffin, petroleum, and tar dissolve in gasoline (petrol), and in benzene (C 6 H 6 ). Cotton and pure paper (like filter paper) will dissolve in strong sulphuric acid. Alcohol (C2H 6 O) dissolves shellac (to make varnish). Again, water will not dissolve carbon disulphide, chloroform, carbon tetrachloride, gasoline or benzene. But it dissolves alcohol in any amount, and ether in limited quantity. Some organic substances, like sugar, dissolve easily in water, but hardly at all in the other solvents just mentioned. Hence candy or molasses can be taken out of cloth by water, but not by solvents for fats. Chemists have as yet found no explanation of these striking differences in solubility. Saturation. As a rule, not more than a certain amount of a solute is dissolved by a given quantity of the solvent. By SOLUTION 109 shaking the solute with the solvent for a sufficient length of time, this maximum amount will finally be dissolved. The solvent is then said to be saturated by the solute in question. Thus, 100 c.c. of water at the ordinary temperature will dissolve as much as 5 g. of potassium chlorate, but not more. The 100 c.c. of water will dissolve 90 g. of Epsom salt, however, before the solvent becomes saturated. On the other hand, a saturated solution of chalk (calcium carbonate) in water will contain only 0.00013 g. in 100 c.c. To describe these cases we should say that potassium chlorate is only moderately soluble in water, Epsom salt very soluble, and chalk insoluble. But no substance is absolutely insoluble. The number of grams of the solute required to saturate 100 c.c. of the solvent we call the solubility of the substance (at the exist- ing temperature). Examples are given below. In some cases there is no limit to the solubility, and therefore no possibility of the solution reaching saturation. Thus alcohol or glycerine and water will mix in any proportion. Such pairs of substances are said to be miscible in all proportions. Dilute and Concentrated Solutions. A dilute solution is one containing little dissolved matter, whether the matter is naturally very soluble or not. A concentrated solution is one containing much of the dissolved substance, and such a solution can be made with very soluble solutes only. Temperature and Solubility. The solubility of every substance in any solvent varies more or less with the temperature. The solubility of niter (potassium nitrate KNOs) in water shows great variation, namely 13 g. in 100 c.c. at 0, 26 g. at 20, 140 g. at 70. On the other hand, the solubility of common salt (sodium chloride NaCl) is more constant, 35.5 g. at 0, 36.5 g. at 20, 38 g. at 70, 40 g. at 100. Usually, as in these two cases, the solubility increases with 110 TEXT BOOK OF CHEMISTRY rise in temperature, but often it diminishes. Thus, the solu- bility of slaked lime (calcium hydroxide Ca(OH) 2 , used to make lime water) is 0.175 g. at 20 and 0.079 g. at 100, and that of anhydrous sodium sulphate (Na 2 S0 4 ), 55 g. at 32.5 and 42 g. at 100. The solubility of gases usually diminishes with rising tem- perature. This may be shown by heating cold tap-water in a beaker. The dissolved gases, originally obtained from the air, appear in bubbles on the bottom and sides as the temperature rises. Crystallization. If the solvent has been saturated while warm, and substance is one that is less soluble at lower temper- atures, then, when the temperature falls, the solute begins to come out of solution. The amount appearing, of course, is only the excess beyond what is needed to saturate the solvent at the lower temperature. If the solute is liquid at the new temperature, it appears at first as a cloud of drops, rendering the liquid milky. This may be shown by cooling a solution of phenol (carbolic acid) in hot water. If the solute is solid, then the particles, as they appear, take the form of crystals (Figs. 39-44). These grow by taking on more of the separating solute. If the cooling goes on slowly, very large crystals can finally be obtained. On the other hand, with rapid cooling, new crystals are continually formed, and a fine crystal-meal falls to the bottom of the solution. The crystals in this meal, however, when viewed through a lens, are seen to be just as perfect as the larger ones. When a dilute solution is used, instead of a saturated one, crystals may still be obtained. A part of the solvent must first be removed, however. This may be done, either by boiling the solution for some time, or by leaving it to evaporate in a wide dish in which a large surface is exposed. SOLUTION 111 When the dissolved substance can form a hydrate (p. 68) which is stable at the temperature of crystallization, the crystals are composed of this hydrate. FIG. 39. FIG. 40. FIG. 41. FIG. 42. FIG. 43. Octahedron Square Pris- Rhombic Monosymmetric Asymmetric (Alum) matic (Niter) (Gypsum) (Hydrated cupric sulphate) The whole of the solvent may be boiled off. But in this case, good crystals of the solute are never obtained the residue is usually a crust composed of imperfect crystals. When the substance is more soluble in cold than in hot water, then crystallization is produced by raising the temperature. Crystallization from a Melted Mass. In this connection, it should be noted that there is another way of obtaining crystals. This is to melt the substance (without any solvent), and allow the mass to cool slowly. When a part has solidified, the rest of the liquid is rapidly poured off. Metals and many other fusible sub- stances give good crystals in this way. Water itself, when it freezes, deposits radiating, hexagonal crystals of ice. Supersaturation. When a hot, saturated solution is cooled, there is quite commonly some delay before the crystals begin to appear. The solution, pending the appearance of the crystals, is then said to be supersaturated. In most cases the crystals appear in due course, especially if the liquid is stirred. But certain sub- stances have a tendency to remain indefinitely in the state of supersaturated solution. The hydrates of sodium sulphate (Na2SO 4 , 10H 2 O) and of sodium thiosulphate (photographers' 112 TEXT BOOK OF CHEMISTRY "hypo" Na^Os, 4H 2 0) are of this nature. The addition of a minute crystal of the substance concerned ("inoculation"), how- ever, always starts the crystallization (Fig. 44). Heat of Solution. Most substances absorb heat when they dissolve, making the solvent cooler, and give out heat when they crystallize. Thus, in the two cases last men- tioned, when the crystallization is brought about in the cool, supersaturated solutions, the rise in temperature is considerable. This fact has been utilized in devising a sort of hot-water bottle. They are rubber bags con- taining a supersaturated solution of sodium ace- tate. Whenever the heat is wanted, the stopper is taken out, rubbed with the finger, and screwed back. This spreads on the inner surface of the stopper, next the liquid, some of the crystals adher- ing to the screw, and so starts the crystallization. The bottle then becomes warm. After being thus used, the bag is placed, without being opened, in boiling water to redissolve the crystals, and, when cold, is ready for use again. Influence of the Solute upon the Solvent. The dissolving of a substance alters the properties of the solvent. The observed changes may be divided into two classes. In one of these two classes equal numbers of dissolved mole- cules of different substances produce the same amount of change. The effect appears thus to be due largely to mechanical causes. Of this nature are the lowering in the vapor pressure of the liquid (p. 113), the raising in its boiling-point (p. 114), and the lowering in its freezing-point (p. 114). In the other class, the amount of the change varies with the substance dissolved. Very striking and difficult to explain, for example, are the erratic changes in volume which occur when SOLUTION 113 solution takes place. Effects of this class show that chemical changes often accompany solution. Thus, this cohesion tends to pull the materials into smaller volume than that of the constituents. For example, 58.5 g. of salt (volume 27.5 c.c.) and 10,000 c.c. of water have a volume, totalling 10,027.5 c.c., but, when they are dissolved, the solution measures only 10.016.5 c.c. This is a very dilute solution (about \ per cent), so that the contraction of 11 c.c. is relatively considerable. To appreciate the real significance of this diminution in volume, we must remember that water and solids require enormous pressures to diminish their volumes. To reduce the volume of 10,000 c.c. of water by 11 c.c., for ex- ample, requires a pressure of 22 atmospheres, or 330 pounds per square inch. In a 20 per cent solution of salt, the shrinkage of 26.7 c.c. per liter corresponds to the effect of about twenty-four times the above-mentioned pressure upon water. The salt and water, therefore, must be pulled together by a truly enormous force. On the other hand, 214 g. of ammonium chloride (volume 142.5 c.c.) and 843.5 c.c. of water have a total volume of 986 c.c., but when dissolved give 1000 c.c. of solution. Here there is an expansion of 14 c.c. In the case of table-sugar and water, how- ever, there is almost no change in volume. Vapor Pressure of Solutions. This cohesion, by holding in the molecules of the solvent, tends to diminish the vapor pres- sure of the latter. Hence aqueous solutions have lower vapor pres- sures than has the pure solvent (water). If the substance is very soluble, and the lowering in the vapor pressure very great, the solution may give a vapor pressure of water less than that commonly present in the atmosphere. Such a solution, placed in an open vessel, will not evaporate. On the contrary, it will take up moisture from the air and increase in bulk. For this reason very soluble substances are commonly moist and, when exposed to the air, extract water from the latter and dissolve in this water. This behavior is called deliquescence, 114 TEXT BOOK OF CHEMISTRY and is shown, for example, by the hydrate of calcium chloride CaCl 2 , 6H 2 0, used to dry gases (p. 60). Boiling-Points of Solutions. Since this cohesion restrains the liberation of vapor, it naturally makes the solution more difficult to boil than the pure solvent. It raises the boiling-point to a higher temperature. It was shown by Raoult (1886), that, in dilute solutions, equal numbers of molecules of different solutes raise the boiling-point of a given solvent to the same extent.* Thus, one molecular weight of sugar (Ci 2 H 22 Oii = 342 g.) or of glycerine (C 3 H 8 3 = 92 g.), dissolved in 1000 c.c. (1 liter) of water, will each raise the boiling- point from 100 to 100.52. Molecular weights of non-volatile substances can be measured by simply finding out what weight of the substance, when dissolved in 1000 c.c., is required to raise the boiling-point of water from 100 to 100.52. Freezing-Points of Solutions. The cohesion between the solvent and solute similarly tends to prevent the freezing of the solution, for freezing means the separation of a part of the pure solvent in the form of ice. Hence solutions can be frozen only at temperatures below those of the pure solvents. Thus, one molecular weight of a substance,* such as sugar (342 g.) or of glycerine (92 g.), dissolved in 1000 c.c. of water, will cause the water to freeze at 1.86 instead of 0. Molecular weights can be measured by this method also. This behavior explains why sea water is much less often frozen in cold weather than fresh water. It explains also why salt thrown on ice will cause the latter to melt. Saturated salt solution freezes only at 21 ( 6 F.), to give a mixture of pure ice and pure salt, both in solid form. Hence, ice and salt can not permanently exist together above that tem- * For important exceptions, see p. 168. SOLUTION 115 perature. When the outside temperature is below 6 F., salt will no longer melt the ice. But calcium chloride, which is more soluble, will do so. A mixture of ice and salt, giving the tem- perature 21, is called a freezing mixture. Such a mixture is used in freezing ice cream and ices. Units Used in Expressing Concentrations. The concen- trations of solutions, saturated and otherwise, are sometimes expressed in physical, and sometimes in chemical, units of weight. When physical units are employed, we give the number of grams of the solute held in solution by one hundred grams of the solvent. The solubilities at 18 of one hundred and forty-two bases and salts are given in a table printed inside the cover, at the front of this book. When chemical units of weight are employed, two different plans are possible, and both are in use. Either the equivalent (p. 53) or the atomic weights may be taken as a basis of measurement. In the former case, the solutions are called normal solutions, and in the latter, molar solutions. A normal solution contains one gram-equivalent of the solute in one liter of solution (not in 1 1. of solvent) . The word "equivalent " has been used hitherto only of elements, and this application of the expression involves an extension of its meaning. An equivalent weight of a compound is that amount of it which will interact with one equivalent of an element. Thus, a formula-weight of hydro- chloric acid HC1 (36.5 g.) is also an equivalent weight, for it con- tains 1 g. of hydrogen, and this amount of hydrogen is displaceable by one equivalent weight of a metal. A formula-weight of sulphuric acid H 2 SO4 (98 g.), however, contains two equivalents of the compound, and a formula-weight of aluminium chloride A1C1 3 (133.5 g.) three equivalents. Hence normal solutions of these three substances contain, respectively, 36.5 g. HC1, 49 g. H 2 SO 4 , and 44.5 g. A1C1 3 per liter of solution. The special property of normal solutions is, obviously, that equal volumes of 116 TEXT BOOK OF CHEMISTRY two of them contain the exact proportions of the solutes which are required for complete interaction. Solutions of this kind are much used in quantitative analysis. We frequently use also decinormal or one-tenth normal solutions (0.1 N or JV/10), and seminormal (0.5 N or N/2), and six times normal solutions (6 N), and EO forth. A molar solution contains one mole (gram-molecular weight) of the solute in one liter of solution (not in 1 1. of solvent) . When molec- ular formulae (p. 77) are used, this means one gram-formula weight per liter. In the cases cited above, the molar solution contains 36.5 g. HC1, 98 g. H 2 S0 4 , and 133.5 g. A1C1 3 per liter. As will be seen, the concentrations of molar and normal solutions are necessarily identical when the radicals are univalent. Definition of a Saturated Solution: A Warning. To avoid a common misconception, it must be noted that a saturated solution must not be defined as one containing all of the solute that it can hold. A supersaturated solution holds more. The saturated solution is one which contains all of the dissolved solute that it can acquire from the undissolved solute. Better still, it is that solution which, when placed with excess of the solute, is found to be in equilibrium. It must be clearly understood that solution is not a process of filling the pores of the liquid. If that were true, approximately equal weights of all substances would find accommodation in equal volumes of water. The fact is that, for example, 100 c.c. of water can dissolve 195 g. of silver fluoride, but only 0.00000035 g. of silver iodide, although the space available (if there is any free space) is the same in both cases. The same conclusion is reached when we consider that two forms of the same compound have different solubilities. Thus, at 20, NauSCX, 10H 2 can give about 18 g. of Na 2 SO 4 to 100 c.c. of water. But anhydrous sodium sulphate Na 2 S0 4 at 20 gives 59 g. to the same amount of water. SOLUTION 117 Is Dissolving a Physical or a Chemical Change ? This is a question still much discussed amongst chemists. Prob- ably in simple, typical cases, like dissolving paraffin in gasoline or benzene, the cohesion may be considered purely physical, and the solution contains both components in unchanged chemical condition. On the other hand, when water is used, as it is more frequently than any other solvent, chemical changes undoubtedly take place. The water itself, at least, is always changed. It is composed chiefly of molecules of dihydrol (H 2 0)2, with some trihydrol (H 2 0) 3 near to 0, and increasing amounts of monohydrol H 2 at higher temperatures. Dissolving any substance in water must upset the equilibrium amongst these different kinds of molecules: (H 2 0) 2 ^2H 2 and produce more of one and less of the other kind. This is the extent of the chemical change in the water. The dissolved substance probably does not always combine with any of the water. But there is evidence that, in many cases, it does. The nature of the compounds is hard to determine, and no simple statement can as yet be made about them. But the compounds, whatever they are, are physically dissolved in the rest of the water. Dissolving, therefore, is partly a chemical, and only partly a pure physical process. Exercises. 1. Give two ways of separating a mixture, con- sisting of a suspended solid and a liquid (p. 107). 2. If you had a spot on your clothing consisting of: (a) grease, (b) sugar, or (c) sugar and grease together, or (d) varnish, how should you proceed in each case to remove the spot ? 3. If chalk (5 g.) and potassium chlorate (5 g.) were mixed, how should you separate them (p. 13) ? Explain how you could secure each substance. 118 TEXT BOOK OF CHEMISTRY 4. Could you make (a) a concentrated, (b) a saturated solution of chalk in water (p. 109) ? Of alcohol in water? 5. If you saturated 200 c.c. of water at 70 with (a) salt, or (b) potassium nitrate, and then cooled the clear liquid to 20, what weight of the solid substance would separate out in each case (p. 109)? 6. To make as concentrated a solution of lime water as possible, should you use hot water or cold (p. 110). 7. Explain why boiled water has a slightly different taste from tap-water that has not been boiled (p. 66). 8. Explain why potassium carbonate becomes wet, and finally dissolves, when exposed to moist air. How must calcium chloride be preserved from becoming moist ? 9. If 52 g. of a substance dissolved in 1000 c.c. of water gives a solution boiling at 100.26, what is the molecular weight of the substance (p. 114)? 10. If 68.5 g. of a substance, dissolved in 500 c.c. of water gives a solution freezing at 1.86, what is the molecular weight of the substance (p. 114)? 11. How much glycerine (C 3 H 6 3 ) could you dissolve in 1 liter of water, and still be able to freeze the water with ice and salt (p. 114)? 12. Explain why a sodium acetate hot-water bottle can be used over and over again. What is the source of the heat it gives out each time it is used ? CHAPTER X HYDROCHLORIC ACID. CALCULATIONS THUS far, the substances we have studied have been mainly air and its components and water and its constituents. Another of the simpler, familiar substances, common salt or sodium chloride (NaCl) may now be taken up. Large amounts of it are used in the household, in cooking and in making freezing mixtures. Still larger quantities are consumed in manufacturing washing soda and soap, for both of which it supplies the necessary sodium. It is employed also to furnish the chlorine for bleaching materials. We shall consider it first as a means of making compounds of chlorine. Preparation of Hydrogen Chloride. When a few drops of commercial, concentrated sulphuric acid (H 2 SO 4 ) are poured upon common salt in an open dish, vigorous effervescence be- gins. This indicates that a gas is forming bubbles upon the salt and that the bubbles are rising through the layer of acid and bursting. The gas is itself invisible, but when we breathe upon the contents of the vessel, a heavy fog is produced. This is due to condensation of water vapor (hi the breath) to droplets of water, in which the gas has dissolved. The fog is composed, in fact, of drops of a solution of hydrogen chloride (HC1) in water, which receives the name of hydrochloric acid (in commerce, muri- atic acid). In order to handle the gas more readily, the sulphuric acid may be allowed to fall from a funnel, drop by drop, upon salt con- tained in a flask (Fig. 45). Soon the air in the flask is all dis- placed by the gas, and the latter issues from the open, delivery 119 120 TEXT BOOK OF CHEMISTRY tube. If a U-tube containing some water is attached to the delivery tube, the gas dissolves in the water as fast as it is formed. If the correct proportions of the materials are used, then, when the action is over, all that remains in the flask is a white solid, different from salt, and called sodium-hydrogen sulphate NaHS0 4 . A part of this may be in solution in a little water, contained originally in the commercial sulphuric acid, of which water commonly forms from 6 to 7 per = ' cent. The equation is easy to make 1 from the formulae given, NaCl + H 2 S0 4 * HC1 T + NaHS0 4 , and requires no further balancing. Other Sources of Hydrogen Chlo- ride. Chlorides of other metals could be substituted for the sodium chloride in this action, and all but the less soluble ones would give hydrogen chloride freely. Com- mon salt is employed because it is the cheapest of the chlorides. While theoretically any acid would, like sulphuric acid, furnish the required hydrogen, and liberate hydrogen chloride, yet in practice no other acid works so well. Some, like phosphoric acid H 3 P0 4 , act too slowly, because they do not dissolve sodium chloride so readily. Others, like hydrofluoric acid HF, are too volatile, and the heat of the action would send them over with the hydrogen chloride in the form of vapor. Sulphuric acid, itself, does not boil below 330. Others, like nitric acid HNOs, would react chemically with hydrogen chloride. Still others, like hydriodic acid HI, being gases, could be used only in aqueous solution, and the water would dissolve the hydrogen chloride produced, and prevent its escape from the vessel. Aside from these objections, all the other acids are more expensive than sul- phuric acid. HYDROCHLORIC ACID 121 The Molecular View of the Interaction of Sulphuric Acid and Salt. One who has used the above-described methods for making hydrogen chloride without reflection would not realize the complexity of the machinery by which the result is achieved. The means are apparently very simple. Yet the mechanical features of this experiment, when laid bare, are extremely curious and in- teresting. A single fact will show the possibilities which are con- cealed in it. If we take a saturated solution of sodium-hydrogen sulphate in water and add to it a concentrated solution of hydrogen chloride in water (concentrated hydrochloric acid), we shall perceive at once the formation of a copious precipitate. This is composed entirely of minute cubes of sodium chloride: NaHS0 4 + HC1 -> H 2 SO 4 + NaCl j .* (2) Now this action is nothing less than the precise reverse of (1), yet it proceeds with equal success. In fact, this chemical interaction is not only reversible (pp. 120, 121), but can be carried to comple- tion in either direction. It is only in presence of a large amount of water that it stops midway in its career and is valueless for securing a complete transformation in either direction: NaHS0 4 + HC1 <= H 2 SO 4 + NaCl. In an action which is reversible, if the products remain as per- fectly mixed and accessible to each other as were the initial sub- stances, their interaction will continually undo a part of the work of the forward direction of the change. Hence, in such a case the reaction must, and does, come to a standstill while as yet only partly accomplished; but this was not the case with actions (1) and (2). Let us examine the means by which the premature cessation of each was avoided. In equation (1) the salt dissolved to some extent in the sulphuric acid, NaCl (solid) <=* NaCl (dslvd.), and so, by contact of the two * See footnote to p. 4. 122 TEXT BOOK OF CHEMISTRY kinds of molecules, the products were formed. On the other hand, the hydrogen chloride, being insoluble in sulphuric acid, escaped as fast as it was formed: HC1 (dslvd.) ^ HC1 (gas). Hence, in that case, almost no reverse action was possible, and the double decom- position went on to completion. With all the sodium-hydrogen sulphate in the bottom of the flask, and most of the hydrogen chloride in the space above, the two products might as well have been in separate vessels so far as any efficient re-interaction was concerned. This plan, in which water is purposely excluded, forms therefore the method of making hydrogen chloride. In equation (2), on the other hand, the hydrogen chloride was taken in aqueous solution, and was mixed with a strong solution of sodium bisulphate. The acid was, therefore, kept permanently in full contact with the sodium bisulphate. It had in this case, every opportunity to interact with the latter and no chance of escape. Every molecule of each ingredient could reach every molecule of the other with equal ease. Furthermore, the sodium chloride, produced as a result of their activity, is not very soluble in con- centrated hydrochloric acid (far less so than in water), and so it came out as a precipitate: NaCl (dslvd.) <= NaCl (solid). But this was almost the same as if it had gone off as a gas. It meant that the greater part of the salt was in the solid form. It was in a state of fine powder, it is true. But, in the molecular point of view, the smallest particle of a powder contains millions of mole- cules, and most of these are necessarily buried in the interior of a particle. Thus, the sodium chloride was no longer able to interact effectively molecule to molecule with the other product, the sul- phuric acid. Hence, there was little reverse action to impede the progress of the primary one. Thus (2) is nearly as perfect a way of liberating sulphuric acid as (1) is of liberating hydrogen chloride. This discussion is given to illustrate the displacement of a chemi- cal equilibrium, and to explain the method of preparing hydrogen chloride. It also throws an interesting light on chemical affinity, HYDROCHLORIC ACID 123 however. Considering action (1), by itself, we might reason that the hydrogen chloride was formed because the affinity of the hydro- gen (H) for chlorine (Cl) was greater than for the sulphate radical (SO 4 ). But, if we did so, then in action (2) we should be compelled to reason similarly that the preponderance of affinity was just the opposite. In point of fact, no conclusion about relative affinity can be drawn from these actions. The effects of affinity are here entirely subordinated by the effects of a purely mechanical ar- rangement, depending on solubility. When the activities of the acids are properly compared, hydrochloric acid is found to be considerably more active than sulphuric acid. Physical properties. Hydrogen chloride is a colorless gas. It is sour in taste, and has a sharp odor. Breathing even a very little causes coughing. It is irritating, but not poisonous in small amounts. The gas is exceedingly soluble in water, one volume of which, at 15, will dissolve no less than 455 volumes of the gas. The saturated solution at 15 contains nearly 43 per cent of the gas by weight. The concentrated hydrochloric acid of commerce con- tains about 35 per cent. The density of the gas (weight of 1 c.c.) is 0.001628. Of more interest to the chemist is the weight of 22,400 c.c. or 22.4 liters (the gram-molecular volume), namely 36.468 grams. This is the molecular weight of the substance. As we have seen (p. 77), it is made up of 1.008 g. of hydrogen combined with 35.46 g. of chlorine. Is the gas heavier or lighter than air? This question is an- swered at once if we recall the fact that the 22.4-liter cube-full of air weighs 28.95 g. (p. 86). The gas is one-fourth heavier. The gas can be liquefied by pressure alone at any temperature below 52 (its critical temperature). One atmosphere pressuie will liquefy it at 84, which is therefore the boiling-point of liquefied hydrogen chloride. 124 TEXT BOOK OF CHEMISTRY When the concentrated aqueous solution is heated, it is the gas and not the water which is driven out, for the most part. When the concentration has been reduced to 20.2 per cent, the rest of the mixture distils unchanged at 110. This occurs because, at this concentration, the gas is carried off in the bubbles of steam in the same proportion in which it is present in the liquid. If a dilute solution is used, water is the chief product of distillation (about 100), but gradually the boiling-point rises and, when the con- centration has reached 20.2 per cent once more, the same hydro- chloric acid of constant boiling-point (110 at 760 mm.), as it is called, forms the residue. The Liquefiability of Gases. It will help us to remember which gases are hard to liquefy and which easy, if we memorize the fact that Faraday (from 1823 to 1845) liquefied most of the familiar gases, and failed only with three, which are the difficult cases, namely hydrogen, oxygen and nitrogen. Nitric oxide NO, and carbon monoxide CO are also difficult to liquefy. The six "inert" gases (p. 253) were not known in his time, and these must now be added to the difficult class. Chemical Properties. In the case of a compound, the chemi- cal property in regard to which we first enquire is its stability. Is it easy or difficult to decompose by heating? Hydrogen chloride can be heated to 1800 before even a trace of it is dis- sociated into hydrogen and chlorine. It is, therefore, very stable. When hydrogen chloride is mixed with ammonia gas, as the latter rises from aqueous ammonia, for example, a direct com- bination takes place. The gases unite to form a cloud of fine, solid particles of ammonium chloride. The properties of the aqueous solution, hydrochloric acid, are es- pecially interesting. The solution is sour in taste. This property is shown by all acids which are soluble in water. HYDROCHLORIC ACID 125 With certain organic compounds it gives new compounds pos- sessing a very characteristic color, and the appearance of this color is used as a test. Thus azolitmin, contained in litmus (an extract from certain lichens), is deep violet-blue in color. When a drop of even very dilute hydrochloric acid is added to litmus solution, the azolitmin changes into a bright red substance. This effect is produced by all soluble acids. Often strips of filter paper, dipped in litmus solution and dried (litmus-paper) are used. When dipped in the acid they turn pink. Certain metals, namely those above hydrogen in the order of activity (p. 54), when placed in hydrochloric acid, displace and liberate hydrogen gas. The same metals, as we have seen (p. 54), liberate hydrogen from all soluble acids. Hydrochloric acid interacts with many compounds in solution. In some instances, as, for example, when the acid is added to a solution of silver nitrate (AgNOs, lunar caustic), one of the new substances produced can be seen, because it appears as a pre- cipitate (see below). The precipitate is silver chloride (AgCl), and is white and curdy in appearance. The other product, nitric acid (HNOs), remains dissolved and invisible: HC1 + AgN0 3 -> AgCl I + HN0 3 . The downward arrow indicates the precipitate. This sort of action, whether it gives a precipitate or not, is always a double decomposition (see below). Such double decompositions are produced by all acids, although the products, unless they are in- soluble, are not visible. If we look back at these four properties of the solution, we see that they are properties common to all acids. If, therefore, we sum up, by saying that in solution hydrogen chloride is an acid, we have given in one word all the important properties of the solution. Precipitation. When two soluble substances, which are capable of acting chemically upon one another, are dissolved, and 126 TEXT BOOK OF CHEMISTRY the solutions are mixed, the action takes place between the dis- solved materials. If one of the products is not soluble, then a supersaturated solution (p. Ill) of this product is at once produced. As a rule, this substance then becomes visible as a fine powder, called a precipitate, suspended in the liquid. Often the precipitated product can be recognized by the physi- cal appearance of the precipitate, and so this sort of action is used as a test for one of the original substances. Thus, precipitates are classified according to their color. Again, precipitates of the same color differ in physical appearance, and may be described as gelatinous, curdy, pulverulent, or crystalline. In the two former cases, at least, the precipitation is so sudden that there is not time for crystals to be formed, and the product is amorphous (Greek, without form) . Double Decomposition. In this chapter we meet for the first time with another variety of chemical change. If we ex- amine the equation for the action of silver nitrate on hydrochloric acid (p. 125), we shall see that the silver nitrate decomposed as if it had been made up of two parts, namely (Ag) and (NO 3 ). The hydrochloric acid similarly separated into its two parts (H) and (Cl). The (Ag) then united with the (Cl) and the (H) with the (N0 3 ). (Ag) (NO,) + (H) (Cl) ** (Ag) (Cl) + (H) (NO,). Since both original substances decomposed, this whole change is called a double decomposition. A sort of exchange between the halves of the decomposing substances took place. The hydrogen chloride was prepared by an action (p. 120), which, if we write it as follows, is seen to be of the same class : (Na) (Cl) + (H) (HS0 4 ) - (H) (Cl) + (Na) (HS0 4 ). The Varieties of Chemical Change. Almost all chemical changes belong to one or other of the varieties we have already HYDROCHLORIC ACID 127 met with and defined (pp. 14, 16, 51). These, along with one example of each, are now placed together: 1. Combination: Zn + S - ZnS. 2. Decomposition: 2KC10 3 -> 2KC1 + 30 2 . 3. Displacement: Zn + H 2 SO 4 -> H 2 + ZnSO 4 . 4. Double Decomposition: AgN0 3 + HC1 > AgCl + HN0 3 . In the first, 2 substances give 1 substance. In the second, 1 substance gives 2 (or more) substances. In the third, 1 element and 1 compound give 1 element and 1 com- pound. In the fourth, 2 compounds give 2 compounds. 5. The actions not included above are those where one sub- stance gives one (different) substance. This kind of change is called internal rearrangement. We shall meet with very few examples of it. This classification suffices for most purposes. But, for special kinds of cases, some other names are used. Thus, a dissociation (p. 68) is an action which belongs to both of the first two classes, be- cause it is reversible. For example, 2H 2 * 2H 2 + 2 . Again oxidation (p. 101) and reduction (p. 101) are connected with the particular substances, such as oxygen, which are con- cerned in the action. The first classification (Nos. 1 to 5) paid no attention to the kinds of elements which were present. Thus, every decomposition is a decomposition. If it is reversible, then it is also a dissociation. If oxygen is set free, then it is a reduc- tion as well. Uses of Hydrochloric Acid. This substance is used as a source of chlorine. It is employed for cleaning metals. Al- though present in very small proportions (about 1 part in 500) in the gastric juice of the stomach, it is a most important component of this fluid. It is sometimes given as a medicine, when the nat- ural supply is too small. 128 TEXT BOOK OF CHEMISTRY CALCULATIONS Calculations Connected with Formulae. In a previous chapter (p. 100) we have seen that formulae represent the com- position of substances, and we have seen how the formula of each substance is worked out from the data obtained by experi- ment. Some ways in which the information contained in formulae can be used may now be taken up. The Composition from the Formula. Take, for example, the formula for silver nitrate, AgNO 3 . To learn the composition of this compound, we look up the atomic weights (see rear cover of this book). We find Ag = 107.88 parts of silver, N = 14.0 parts of nitrogen, Os = 3 X 16.0 or 48.0 parts of oxygen. The proportions of the constituents, in the same order, therefore, are 107.82 : 14 : 48. What is the proportion of nitrogen to oxygen alone? It is 14 : 48, or, dividing by 14, it is 1 : 3.43. Significant Figures. The division of 14 into 48 really gives the quotient 3.428571. But no atomic weights have been meas- ured so accurately that we know the values of the numbers be- yond the second place of decimals. Hence 3.43 is just as likely to be the exact value as the longer number. In the most exact cal- culation we round the number off at the second decimal. For rougher purposes we round it off at the first decimal. The Formula-Weight. The sum of the weights of the constituents indicated in the formula is called the formula-weight. For silver nitrate this has the value, 107.88 + 14.0 + 48.0, or 169.88. If the substance is a gas, or is easily volatile, the formula-weight will be also the molecular weight. Thus, acetylene gas (used in lighting) has the formula C 2 H 2 . The composition is C 2 = 2 X 12 or 24 parts of carbon and H 2 = 2 X 1.008 or 2.016 parts of hydro- CALCULATIONS 129 gen. The molecular weight is 24 + 2.016 = 26.016. Again, hydrogen peroxide (used in medicine), a liquid, is decomposed when boiled, but it dissolves in water and depresses the freezing- point. Its molecular weight has therefore been determined (p. 114) and is H 2 O2. The molecular weight on which this formula is based is 2.016 + 32, or 34.016. When the substance is neither volatile nor soluble, the simplest formula is always used, and therefore only the formula-weight can be ascertained. To Find the Percentage Composition. In silver nitrate the proportions are 107.88 of silver, 14.0 of nitrogen, and 48.0 of oxygen in a total of 169.88. In one hundred parts, the silver will 107 R8 14- 16988 ' - 16988 the oxygen X 100 = 28.25. The same results may be obtained by the rule of proportion. Thus, for the silver, 169.88 : 107.88 :: 100 : x where x is the per- centage of silver. Calculations by Use of Equations. We frequently desire to know what weight of a product can be obtained from a given weight of the necessary materials. For example, what weight of zinc sulphide can be made with 100 g. of sulphur? It is under- stood, of course, that the necessary zinc is available. In such calculations mistakes are easily made. The follow- ing rules must be strictly followed : First, write down the equation: Zn -f S - ZnS. Second, place beneath each formula the weight for which it stands: Zn + S > ZnS. 65.37 32.07 97.44. 130 TEXT BOOK OF CHEMISTRY Third, read the whole statement. In this case it reads: 65.37 parts by weight of zinc combine with 32.07 parts of sulphur to give 97.44 parts of zinc sulphide. Fourth, re-read the original problem. Then place the amount given in the problem (100 g. of sulphur) under the formula of the substance concerned. Then observe that the problem asks "What weight of zinc sulphide?" and place an interrogation point or an x under the formula of that substance: Zn + S -> ZnS 65.37 32.07 97.44 100 ? or x Fifth, read the problem as it now appears in this expanded equation: 32.07 g. of sulphur will give 97.44 g. of zinc sulphide, therefore 100 g. of sulphur will give x g. of zinc sulphate. Sixth. The answer may be now obtained by stating the pro- portion in the same order : 32.07 : 97.44 :: 100 : x ( = 303.8). If the expanded equation has been prepared correctly, this final statement is purely mechanical. It will be seen that only two of the three quantities in the equation were really used. Alternative to the Sixth Step. We may also say: If 32.07 g. of sulphur will give 97.44 g. of zinc sulphide, 1 g. of sulphur 97 44 will give ' g. ( = 3.038 g.) of zinc sulphide. Then, if 1 g. of 32.07 sulphur gives 3.038 g. of zinc sulphide, the 100 g. of sulphur will give 100 X 3.038 g. ( = 303.8 g.) of zinc sulphide. Warnings. In solving the exercises at the end of the chapter, beware of three kinds of mistakes commonly made by beginners. 1. Conquer a tendency to say that the symbols Zn and S stand for " 1 part" of zinc or of sulphur. They stand for 1 chemi- cal unit, or atomic weight, or atom, in each case, that is to say, for 65.37 " parts" and 32.07 "parts," respectively. CALCULATIONS 131 2. Follow the rules laid down above. When one has once become familiar with the art of solving such problems, running through the rules takes only a few seconds. The chemist does it almost unconsciously. The beginner always thinks he can ignore these rules, and he fails in consequence. Writing the equation in the expanded form, and then reading the problem into it are absolutely essential steps. 3. Do not read the original problem carelessly and make the equation backwards, that is, with the sides reversed. If there seems to be confusion somewhere, when the last steps are reached, this hint will probably show the cause of the difficulty. Another Example. What weight of hydrogen is required to reduce 45 g. of magnetic oxide of iron to metallic iron ? Following the rules, as before, we reach the expanded equation: Fe 3 4 3 X 55.84 + 4 X 16 167.52 + 64 231.52 45 g. x Observe that the atomic weights are multiplied by the sub-num- bers, so that, for example, Fes = 3 X 55.84. Observe also that the formula weights are multiplied by the coefficients, when such occur, in front of the formula, so that, for example, 4H 2 = 4 X 18.016. The proportion 231.52 : 8.064 :: 45 : x (= 1.57) supplies the answer, 1.57 grams of hydrogen. Using the alternative plan (p. 130): If 321.52 g. of magnetic oxide are reduced by 8.064 g. of hydrogen, 1 g. will be reduced o 0^4 by ' g. (= 0.035 g.) of hydrogen. Hence, if 1 g. of magnetic 2ol.52 oxide is reduced by 0.035 g. of hydrogen, 45 g. will be reduced by 45 X 0.035 g. (= 1.57 g.) of hydrogen. 4H 2 -> 3Fe H h 4H 2 0. 8 x 1.008 3 X 55.84 4(2 X 1.008 + 16) 8.064 167.52 4 X 18.016 8.064 167.52 72.064 132 TEXT BOOK OF CHEMISTRY Exercises. 1. Complete the equation ZnCl 2 + H 2 S0 4 > ZnSC>4 + , and attach the name of the substance to each for- mula in it. 2. Point out the differences in physical properties between oxygen and hydrogen chloride. 3. Make equations for the displacement of hydrogen from hydrochloric acid by zinc and by sodium (pp. 52 and 53). 4. Give additional examples of the first four varieties of chemi- cal change (p. 127). 5. Classify (p. 127) the following actions: (a) the action of steam on iron (p. 51); (b) the rusting of iron; (c) the electrolysis of dilute hydrochloric acid (p. 55); (d) the effect of heating the hydrate of cupric sulphate (p. 111). 6. Is the decomposition of mercuric oxide a case of: (a) disso- ciation, or (b) oxidation or reduction? 7. What is the proportion of: (a) sodium to one part of chlorine in salt; (b) one part of hydrogen to nitrogen in nitric acid (p. 315). 8. Calculate the percentage composition of: (a) sulphuric acid; (b) acetylene C2H 2 . 9. What weight of hydrogen is displaced by the action of 100 g. of zinc upon an excess of hydrochloric acid (ZnCl2 is formed) ? 10. What weight of silver chloride is formed by the interaction of 5 g. of sodium chloride with silver nitrate? 11. What should be the difference in cost of 1000 g. of oxygen according as it is obtained by decomposing mercuric oxide at $3.00 (12/6) per kilog. or potassium chlorate at 30 cents (1/3) per kilog? CHAPTER XI CHLORINE. CALCULATIONS CHLORINE was discovered by Scheele (1774). It was supposed to be a compound containing oxygen, until the contrary was proved by Davy (1809-1818). It is used in immense quantities for making bleaching agents. Occurrence. Many compounds of chlorine occur in nature, but the most plentiful is common salt (NaCl). Most of the mat- ter dissolved in sea water is sodium chloride. Salt also occurs underground, either in strata in almost pure form, or mixed with rocky material. Near such deposits, wells and springs of salt water are common. Preparation by Electrolysis of a Chloride. Chlorine is liberated by passing a current of electricity through an aqueous solution of a chloride, such as hydrogen chloride or sodium chloride (see Fig. 27, p. 55). Much of it is, in fact, manufactured by electrolyzing natural brines. The chlorine is liberated, first in solution and later as a gas, at the positive wire (anode). Other Products of Electrolysis of Chlorides. In elec- trolyzing hydrochloric acid, the hydrogen is set free at the nega- tive wire (cathode). With sodium chloride we might expect to get free sodium. It will be recalled, however, that sodium is very much more active than is hydrogen, and indeed displaces hydrogen from water. Hence the electrical energy sets free the more easily liberated element the hydrogen and the sodium remains in 134 TEXT BOOK OF CHEMISTRY the solution as sodium hydroxide (NaOH). The process is best shown by a diagram: j2Na;Cl -> C1 2 T (neg. wire) (pos. wire) T H 2 ^-2HJOHj The chlorine and hydrogen H 2 being liberated, leave behind in the solution the constituents of 2NaOH: 2NaCl + 2H 2 O + Elect. -> H 2 + C1 2 + 2NaOH. Preparation from Oxygen and Hydrogen Chloride. When hydrogen chloride and oxygen gases are heated, they inter- act very slowly to give water and chlorine. The action is greatly hastened by contact with cupric chloride. Lumps of pumice, saturated with a solution of ^>H- ,_^ftah*8sgaa^^ Hh^ this catalyst, or contact agent, are placed in a tube. When rIG. 46. the mixture of gases is passed over the heated pumice (Fig. 46), steam, chlorine, and about 20 per cent of unchanged oxygen and hydrogen chloride issue at the other end: Skeleton equation: HC1 + 2 =* H 2 O -f C1 2 . Balanced equation: 4HC1 + O 2 <= 2H 2 O + 2C1 2 (80 per cent). Longer heating does not alter the proportion of the materials successfully transformed. This is Deacon's process. That 80 per cent is changed, and 20 per cent unchanged, is due to the fact that the action is reversible. If we lead pure chlorine and steam through the tube (read the equation back- wards), 20 per cent of hydrogen chloride and oxygen are formed. No more than 20 per cent is formed, because these products are continually being used up again and reproduce steam and chlorine. If one product could be separated (p. 120) from the other, to pre- vent the backward action, the yield could be raised to 100 per cento But all the four substances are gases (in the hot tube), and mix perfectly. CHLORINE 135 The result is interesting, because it shows that, under these conditions, oxygen is more active than chlorine in combining with hydrogen. We can try the experiment by mixing hydrogen with quantities of oxygen and of chlorine, each sufficient to com- bine with all the hydrogen. When we explode the whole mixture, we find that 80 per cent of the oxygen succeeds in uniting with hydrogen, but only 20 per cent of the chlorine is so used. Preparation front Hydrochloric Acid and an Oxidizing Agent. The best way to make a supply of chlorine in the laboratory is to place potassium permanganate crystals (KMnCU) in a flask (Fig. 47) and allow concentrated hydrochloric acid, previously mixed with an equal volume of water, to fall upon them drop by drop. The gas is rather soluble in water, and is best collected by displacing the air from bottles. When one bottle is full, it is stoppered and a fresh one substituted. To avoid the escape of the very irritating gas into the room the tube from the collecting bottle dips beneath sodium hydroxide solution. 2KMnO 4 + 16HC1 -> 8H 2 O + 2KC1 + 2MnCl 2 + 5C1 2 . The essential feature of the chemical change is that the oxygen of the potassium permanganate unites with the hydrogen of the hydrochloric acid to give water. The potassium and manganese take as much chlorine as they require to form their chlorides, KC1 and MnCl 2 . The rest of the chlorine is liberated. This action is considered to be an oxidation of the hydrogen chloride by the permanganate, although the latter only removes the hydrogen, and does not reach the stage of adding any oxygen FIG. 47. 136 TEXT BOOK OF CHEMISTRY to the chlorine. The potassium permanganate, which carried the oxygen, is called the oxidizing agent. Since the permanganate lost oxygen, it was itself reduced. In all oxidations one substance is oxidized and another reduced. Deacon's process (p. 134) is also an oxidation of hydrogen chlo- ride (by free oxygen). The oxygen is reduced to water. Chlorine may be prepared by using other substances to oxidize hydrochloric acid. Amongst those which are suitable are man- ganese dioxide MnC>2, potassium chlorate KClOs and red lead Pb 3 4 . Manganese Dioxide and Hydrogen Chloride. The action of manganese dioxide upon hydrochloric acid is an instructive one. It is a general rule, of which we shall meet many applications, that when an acid interacts with an oxide of a metal, there are two con- stant features in the result, namely: (1) The oxygen of the oxide combines with the hydrogen of the acid to form water, and (2) the metal of the oxide combines with the acid radical of the acid accord- ing to the valences of each. Here the skeleton equation should be MnO 2 + HC1 - H 2 O + MnCl 4 . With O 2 , to form water, 4HC1 is required, and the product is 2H 2 0. Hence the equation is Balanced: Mn0 2 + 4HC1 -* 2H 2 O + MnOU. This is what happens in the first place. The products actually obtained, however, are water, manganous chloride MnCl 2 and chlorine. The manganese tetrachloride can be preserved by cool- ing the mixture. It is decomposed by the heating, the chlorine escapes, and the other two products remain in the vessel. Mn0 2 + 4HC1 -> 2H 2 O + MnCl 2 + C1 2 . (1) We owe the chlorine to the fact that the tetrachloride is unstable. If we had used manganous oxide MnO, we should have had a double decomposition: MnO + 2HC1 -> H 2 + MnCl 2 , (2) CHLORINE 137 but we should have got no chlorine. Perhaps the simplest way to describe the difference between these two actions is in terms of the valence of the manganese. In Mn O2 n the element is quadriva- lent. This means that its atomic weight professes to be able to hold four atomic weights of a univalent element. The four valences of oxygen (2O 11 ) can do the same thing. In equation (1) the oxygen fulfils this promise by taking 4H 1 . But the Mn IV can hold only 2C1 1 , permanently, and lets the other 2C1 T go free. In other words, the valence of the atomic weight of manganese changes in the course of the action. In equation (2), on the other hand, the manganese is bivalent to start with (Mn^O"), and is able to retain the amount of chlorine (2C1 1 ) equivalent to O n . Actions like that of manganese dioxide in (1) are classed as oxidations. The hydro- gen chloride, or rather half of it, is oxidized. A graphic mode of writing may make this remark clearer: , O + 2HC1 -> H 2 + Mn n Cl 2 Mn-f; ^O + 2HC1 -* H 2 O + C1 2 The upper half is a double decomposition, the lower an oxidation by half the combined oxygen of the dioxide. The same explana- tion applies to the interaction of lead dioxide with hydrochloric acid. Physical Properties. Chlorine is a greenish-yellow gas, and takes its name from the Greek for this color. It has an exceed- ingly disgreeable odor and irritates the lining of the nose and throat. Alcohol vapor or ammonia, when breathed, relieve the irritation. The density of the gas is recorded in the formula C1 2 . The 22.4-liter-cube-full weighs 70.92 g., against 28.95 g. for air, so that chlorine is about 2J times heavier. Two volumes of the gas dissolve in one volume of water at 20. The solution is called chlorine-water. 138 TEXT BOOK OF CHEMISTRY The gas is liquefied by pressure below 146 (its critical tem- perature), the pressure required at 20 being 6.6 atmospheres. The liquid boils at 33, and solidifies at 102 (its melting- point) . Chemical Properties. Chlorine is an element with about the same degree of activity as oxygen, and it unites with very much the same list of other elements. The compounds are called chlorides. Unites with Metals. \ When powdered antimony (cold) and iron powder (warmed) are thrown into chlorine, they combine, and red hot particles of the chlorides, SbCl 3 and FeCl 3 , fall to the bottom. Copper leaf (Dutch metal, used in " gilding"), or heated copper foil, burns in the gas, giving a fog of solid cupric chloride CuCl 2 . Skeleton: Sb + C1 2 -> SbCl 3 . Balanced: 2Sb + 3C1 2 7? 2SbCl 3 . Sodium burns brilliantly in chlorine, giving sodium chloride. That a shining metal and a poisonous irritant like chlorine, in uniting, should yield a mild, household article like common salt illustrates very well the extraordinary nature of chemical change. When thoroughly freed from moisture, chlorine no longer combines with metals like copper and iron. Water seems to be needed as a contact agent, in these, as well as in hundreds of other chemical actions. Hence, carefully dried chlorine, in liquid form, can be, and is stored and sold in iron cylinders (see detinning, p. 487). Unites with Hydrogen. A jet of burning hydrogen, lowered into a bottle of chlorine, continues to burn, giving hydrogen chloride HC1, the presence of which is shown by the fog produced by allowing the gas to come in contact with moist air: H, + C1 S ->2HC1. CHLORINE 139 Hydrogen and chlorine, mixed, do not combine when cold, provided strong light is excluded. But sunlight, or light from burning magnesium (" flashlight powder "), starts the combina- tion, which occurs with explosive violence. Plunging a lighted taper into the mixture has, of course, the same effect. Acts upon Compounds Containing Hydrogen. Because of its activity toward hydrogen, chlorine removes hydrogen from many compounds. Thus, if a lighted wax taper be plunged into chlorine, it continues to burn, though with a feebler flame. Dense smoke, composed of particles of free carbon, rise from the flame. Blowing the breath into the jar, afterwards, gives the fog due to hydrogen chloride. Thus the presence of hydrogen and carbon in the wax of the taper is proved. From this we learn, also, that chlorine has a relatively small tendency to combine with carbon. A few drops of warm turpentine (Ci Hi 6 ) upon a slip of filter paper will blaze up in chlorine, giving hydrogen chloride and an immense cloud of soot: Skeleton: Ci H 16 + Cli-> C -f HC1. Balanced: CioH 16 + 8C1 2 -> IOC + 16HC1. Acts upon Water. We have seen that chlorine acts on steam, reversing Deacon's reaction to the extent of 20 per cent. It acts also upon cold water, when dissolved in the latter, although in a similarly incomplete way. With half-saturated chlorine- water at 10, about one-third of the chlorine is used. One of the products, hypochlorous acid HC1O, is of especial interest, because it is an exceedingly active substance, much used as an oxidizing agent (see p. 139) and for bleaching: C1 2 + H 2 <=> HC1 + HC10 (33%). Unites with Non-metals. Phosphorus burns in the gas, giving the vapor of phosphorus trichloride PCls. Sulphur, when heated, combines more slowly, giving sulphur monochloride S 2 Cl2, used in vulcanizing rubber. 140 TEXT BOOK OF CHEMISTRY Although chlorine does not combine very readily with carbon, several compounds of carbon and chlorine are very important, such as chloroform CHC1 3 and carbon tetrachloride CCU (p. 192). Composition of Hydrogen Chlo- ride. Now that we are familiar with the properties of chlorine, as well as with those of hydrogen, we may re- turn to the question of the proportion by volume in which they are produced by decomposition of hydrogen chloride. When we electrolyze hydrochloric acid in the apparatus shown in Fig. 27 (p. 55), we find that the chlorine dis- solves to a large extent in the liquid, and its true volume as gas is not easily ascertained. The apparatus (Brown- lee's) in Fig. 48 avoids the difficulty by enabling us to saturate the liquid with chlorine before observing the col- The volumes of the two gases are found to be equal. FIG. 48. lee ted gases. FIG. 49. A simpler apparatus (Fig. 49) may be used to show the same fact. The gases are generated, by electrolysis, in the test-tube, CHLORINE 141 pass through the straight tube, driving the air before them, and finally bubble through sodium hydroxide solution. The whole apparatus must be covered with a dark cloth to exclude light, and handled in diffused light. In the course of fifteen minutes, the chlo- rine ceases to be dissolved, and the gases come off in their natural proportions. In half-an-hour more they have filled the tube. The stop cocks are now closed, the tube is set in a tall cylinder containing potassium iodide (KI) solution (Fig. 50) and the lower stop cock is opened. The potassium iodide acts upon the chlorine, giving potassium chloride and free iodine, which is a solid, i FIG. 50. and the liquid rises until it fills half the tube. The remaining gas burns and is hydrogen. The volume of the hydrogen chloride, in relation to the volumes of the constituents, may be learned by using a different apparatus (Fig. 51). A test- tube of heavy glass is filled with dry hydrogen chloride and closed with a rubber stopper greased with vaseline. A little sodium amal- gam (solution of so- dium in mercury) is introduced, and the stopper instantly re- placed. When the contents are shaken FIG. 51. for one or two min- utes, the sodium com- bines with the chlorine and the hydrogen is liberated and re- mains. The mouth of the test-tube is then immersed in a jar of 142 TEXT BOOK OF CHEMISTRY water, and the stopper withdrawn. The water rises and fills about half the tube. Conclusion. The hydrogen has half the volume of the hydrogen chloride, and the volume of the chlorine is equal to that of the hydrogen. Therefore: 1 vol. hydrogen + 1 vol. chlorine - 2 vols. hydrogen chloride. This result further illustrates Gay-Lussac's law (p. 60). Confirmation of the Formulse C1 2 and H 2 . According to Avogadro's hypothesis, there are equal numbers of molecules in equal volumes of these gases. Let the squares represent small volumes, containing 1000 molecules each: Hyd. chloride. Hyd. Chlor. [TOGOl pOQ| <- [IOOO| + [lOOOl It appears from this that 2000 molecules of hydrogen chloride come from 1000 molecules of hydrogen and 1000 molecules of chlorine. Now, each molecule of hydrogen chloride contains hydrogen, so that the 1000 molecules of hydrogen must have given 2000 atoms of hydrogen, one for each molecule of the compound. Hence each molecule of hydrogen contains two atoms. The same is true of each molecule of chlorine. This confirms the conclusion we reached before (p. 75), namely, that the formulae of the free gases are H 2 and Cl2, and that single atoms, of the elements occur only in combination (as in NaCl, HC1, etc.). It would be more correct to reason that each molecule of hydro- gen was split in two, and therefore contains at least two atoms. We find no need, however, of supposing that there are more than two atoms in each molecule. CALCULATIONS Calculations Involving Volumes of Gases. For such cal- culations the simplest method is always to use the gram-molecular volume the 22.4-liter cube. CALCULATIONS 143 Relative Densities of Gases. What are the relative weights of equal volumes of chlorine and hydrogen chloride? The for- mulae record the weights of 22.4 liters: C1 2 = 2 X 35.46 = 70.92, and HC1 = 1.008 + 35.46 = 36.468. The required relation is 70.92 : 36.468, so that chlorine is almost twice as heavy, bulk for bulk, as hydrogen chloride. Density Compared with Air. We have seen (p. 86) that 22.4 liters of air weigh 28.95 g. Is water vapor heavier or lighter, at the same temperature and pressure? The formula, H 2 O, shows that 22.4 liters of steam weigh 2 X 1.008 + 16 = 18.016. Air is heavier, in the ratio of 29 : 18, approximately. The Volume of Gas from Given Weights of Material. What volume of oxygen is obtained by heating 15 g. of potassium chlorate? Follow the rules given (p. 129) for calculating weights: 2KC1O 3 2KC1 + 3O 2 . Weights: 2 (39.1 + 35.46 + 3 x 16) 2 (39.1 + 35.46) 3 x 32 or 245.12 149.12 96 Now, add the volumes in the case of substances which are gases. Here oxygen is the only one. Remembering that O 2 ( = 32 g.) occupies 22.4 liters, the volume of the oxygen is 3 X 22.4 liters, or 67.2 liters. Then add the data given, and the x in their proper places. The x goes under the volume of oxygen: 2KC1O 3 - 2KC1 + 3O 2 . 245.12 g. 149.12 g. 96 g. (or 67.2 liters) 15 g. x The problem now reads thus: If 245.12 g. of potassium chlorate give 67.2 liters of oxygen, 15 g. of the former will give x liters of the latter: 245.12 : 67.2 :: 15 :x(= 4.07), where x, the answer, is 4.07 liters of oxygen. 144 TEXT BOOK OF CHEMISTRY Observe that, when we use the gram-molecular volume, one proportion gives the answer. It is not necessary to make two steps in the calculation, by finding first the weight of the oxygen, and then its volume. Another Example. What weight of copper will combine with 15.2 liters of chlorine (at and 760 mm.) ? Cu + Cla -> CuCl 2 . 63.57 g. 70.92 g. or 22.4 1. 134.49 g. x 15.2 liters The proportion is: 22.4 : 63.57 :: 15.2 :x(= 43.14 g.). At Other than Standard Conditions. If the problem concerns a gas at some temperature and pressure not and 760 mm., then correction must be made as a separate calculation. For example, if the 15.2 liters of chlorine, in the foregoing illus- tration, had been at 15 and 742 mm., then the volume would first have been reduced to and 760 mm. It would serve the purpose equally well to correct the 22.4 liters to 15 and 742 mm. Questions of Volume Alone. When weights are not men- tioned in the question, but volumes only, the calculation is very simple. For example: What are the relative volumes of oxygen and hydrogen chloride used in Deacon's process? 4HC1 + 2 -> 2H 2 O + 2C1 2 Volumes: 4 X 22.4 1 X 22.4 2 X 22.4 2 X 22.4 Molecules: 4122 Remembering that equal numbers of molecules occupy equal volumes, and that 4 HC1 = 4 molecules, 2 = 1 molecule, etc., we perceive that 4 volumes of HC1 will be required for 1 volume of oxygen. Does the volume change during the process ? Yes, 4+1 volumes become 2 + 2, or 5 volumes give 4. CALCULATIONS 145 If we measure the substances at a low temperature before and after the action, so that the steam condenses to water, then the volume shrinks to practically 2 volumes of gas (the chlorine). A gram-molecular volume of steam gives only 18 g. (or 18 c.c.) of water. Warnings. The commonest mistake made in these calcula- tions is to neglect to use, in the equations, the molecular formulae. We must use H 2 , O 2 , C2H 2 (acetylene), etc., (and not H, 0, or CH), because the 22.4 liters hold the weights represented by the molecular formulae. Before calculating the volume (gaseous) of a substance repre- sented in an equation, consider whether it is a volatile substance. Only volumes of gases can be calculated by the rules given above. Exercises. 1. What would be the results of electrolyzing aqueous solutions of: (a) potassium chloride; (b) cupric chloride (cf. p. 54). 2. How should you separate the chlorine and the steam produced by Beacon's process? 3. Make equations showing the interactions with hydrochloric acid of: (a) manganese dioxide; (b) potassium chlorate; (c) red lead. Note that the remark about permanganate (p. 135 line 24) applies here also. The metals form MnCl 2 , KC1, and PbCl 2 , respectively. 4. What would be the pressure in a cylinder of liquid chlorine at 20? 5. Make equations for the union of chlorine with: (a) copper; (b) sodium; (c) iron; (d) phosphorus; (e) sulphur. 6. When plunged into chlorine, a jet of illuminating gas con- tinues to burn. A stream of soot rises from the flame, however, and blowing the breath into the jar, afterwards, produces a fog. What do you infer as to the constituents of illuminating gas? 146 TEXT BOOK OF CHEMISTRY 7. What information is conveyed by the fact that the formula of the chloride of sulphur is written 82012, and not SCI. 8. Make the molecular equation for the action of sodium upon hydrogen chloride (p. 102). Why does not the mercury interact with the latter (cf. p. 54) ? 9. What facts led us, in Chap. VII,' to the conclusion that the molecular weight of chlorine was 70.92 while its atomic weight was 35.46? 10. What are the relative densities (p. 75) of: (a) oxygen and chlorine; (b) hydrogen and hydrogen chloride; (c) hydrogen and air; (d) acetylene C 2 H 2 and air? 11. What volume of hydrogen chloride at and 760 mm. is obtained by the interaction of 65 g. of sodium chloride, and an excess of sulphuric acid (p. 120) ? 12. What weight of zinc is required to make 100 liters of hydro- gen, at 10 and 750 mm., by displacement from hydrochloric acid (p. 102)? 13. What are the relative volumes of the factors and of the products in the interaction between: (a) turpentine vapor and chlorine : (b) oxygen and carbon disulphide vapor (giving S02 and C0 2 )? 14. What are the relative volumes of the products in the de- composition of: (a) mercuric oxide; (b) chlorine monoxide C^O (chlorine and oxygen are formed) ? 15. What are the relative volumes of the volatile substances concerned in the action of water vapor on iron (p. 102) ? 16. Using the relative volumes in which oxygen and hydrogen combine to form steam, prove that each molecule of free oxygen is composed of at least two atoms (cf. p. 142). 17. (a) Calculate the weight of chlorine dissolved by 100 c.c. of water at room temperature (p. 137). (b) In half -saturated chlorine-water, what is the concentration of hypochlorous acid (p. 139)? CHAPTER XII ENERGY AND CHEMICAL CHANGE IN describing chemical changes, the fact that heat was evolved has frequently been mentioned. In several instances a current of electricity has been used to produce chemical change. It is now necessary to collect these scattered facts and classify them for future use. Physical Accompaniments of Chemical Change. When iron and sulphur combined (p. 14), and when iron burned in oxy- gen or copper in chlorine, much heat was developed. On the other hand, the decomposition of mercuric oxide, as was pointed out (p. 15), owed its continuance to the persistent application of heat and ceased as soon as the source of heat was withdrawn. Here, apparently, heat was consumed during the progress of the change, and the chemical action was limited by the amount of heat supplied. The production or consumption of heat may, therefore, be a feature of chemical change. In the burning of iron or magnesium in oxygen, and in the actions of chlorine on copper and turpentine, light was also pro- duced. Conversely, silver chloride can be kept any length of time in the dark, but in sunlight it becomes first bluish and then brown, simultaneously giving off chlorine gas and finally leaving only silver as a fine powder. Silver bromide or iodide, in photographic plates, films, and paper, is changed by light in a similar way, liberating the bromine or iodine. It would appear, therefore, that light may be given out or consumed in connection with chemical change. We have seen (p. 55) that a current of electricity may be em- ployed to decompose hydrochloric acid and other chlorides, and 147 148 TEXT BOOK OF CHEMISTRY the battery, or other source of the current must be kept going or the chemical change stops. The inverse of this is likewise familar. If we place in dilute sulphuric acid a stick of the metal zinc, we find that hydrogen is given off (Fig. 52), that the zinc goes into solution as zinc sulphate (p. 102), and that a large amount of heat is developed. If zinc in fine particles, with much surface, is used, the liquid may even rise spontaneously to the boiling-point. This form of the action produces heat. If, however, we attach the same stick of zinc to a copper wire and, having provided a plate of platinum also connected with a wire, immerse the two simultaneously in the acid (Fig. 53), then a galvanometer, with which the wires are connected, shows at once the passage of a current of electricity round the circuit. Exactly the same chemical change goes on as before. The sole FIG. 52. FIG. 53. difference is that the gas appears to arise from the surface of the platinum. It is easy to show, however, that the platinum by ENERGY AND CHEMICAL CHANGE 149 itself is not acted upon by dilute acids and, in this case, under- goes no change whatever; it serves simply as a suitable conductor for the electricity. Here, then, in place of the heat which the first plan produced, we get an electric current. The arrangement is, in fact, a battery-cell, for a battery is a system in which a chemical action which would otherwise give heat furnishes electricity in- stead. Thus, electrical energy may be consumed or produced in connection with a change in composition. Even violent rubbing in a mortar, in the case of some substances, can effect an appreciable amount of decomposition in a few min- utes. In this way silver chloride can be separated into silver and chlorine, just as by light. It is the mechanical energy which is the agent, and part of it is consumed in producing the change, and only the balance appears as heat. Conversely, the production of mechanical energy, as the result of chemical change, is seen in the behavior of explosives and in the working of our muscles. Thus, mechanical energy may be used up or produced in chemical changes. Summing up our experience, we may state that no change in composition occurs without some accompaniments, such as the production or consumption of heat, light, electrical energy, or, in some cases, mechanical energy. Classification of the Accompaniments of Change in Com- position: Energy. The problem of classifying (i.e., placing in a suitable category) things like heat, light, and electricity has occu- pied much attention. In all changes in composition, one of these natural accompaniments is given out or absorbed, sometimes in great amount, yet in none is any alteration in weight observed.* There are many things which are real, however, even if they are not affected by gravitation. In the present instance we reason as follows: * Electrons (q.v.) do possess mass, but it is very small compared with that of the materials concerned. 150 TEXT BOOK OF CHEMISTRY A brick in motion is different from a brick at rest. The former can do some things that the latter cannot. Furthermore, we can easily make a distinction in our minds. The brick can be deprived of the motion and be endowed with it again. Thus, we can get the idea of motion as a separate conception. Similarly, we observe that a piece of iron behaves differently when hot, and when cold, when bearing a current of electricity, and when bearing none. We conceive then of the brick or the iron as having a certain amount and kind of matter which is unalterable, and as having motion, heat, or electricity added to this or removed. Thus, we describe our observations, by using two categories, one of which includes the various kind of matter, and the other, various things whose association with matter, seems to be invariable and is often so conspicuous. The latter we call the forms of energy. The Practical Importance of Energy in Chemistry. The absorption or liberation of energy accompanying a chemical transformation of matter is often, of the two, the more important feature. We do not burn coal in order to manufacture carbon dioxide gas. We are glad to get rid of the material product through the chimney. It is the heat we want. We do not buy gasoline (petrol) for an automobile in order to obtain various gases to ex- pel through the muffler. We really pay for the mechanical energy. It is the same with burning illuminating-gas or magnesium powder when we want light, and with eating food, which we do, chiefly, to get energy to sustain our activity. We do not run electricity for hours into a storage battery, in order to make a particular com- pound (lead dioxide, for example), but in order to save and store the energy for future use. In industry and life fully half the total amount of chemical change involved is set in motion by us, solely on account of the energy changes it involves. But the production of energy in chemical change is not only thus of practical importance; it is also of scientific interest, as will be seen in the section on energy and chemical activity (see below) . ENERGY AND CHEMICAL CHANGE 151 Interconvertibility of Forms of Energy: Conservation. At first sight, these accompaniments of matter seem to be quite unrelated. But a relation between them can be found. If the heat of a Bunsen flame or of the sun is brought under a hot-air motor (Fig. 54) violent motion results. Again, if the motor is connected with a dynamo, electricity may be generated. Still again, if the current from the dynamo flows through an incandes- cent lamp, heat and light are evolved. Conversely, when motion of the hot-air motor is impeded by a brake, heat appears. When a current of electricity is run through the dynamo, the armature of the latter turns and motion results. But the most signifi- cant facts are still to be mentioned. The heat absorbed by the motor is found to be greater when the machine is permitted to move and do work, than when it is not. Thus, it is found that when work is done some heat disappears, and this heat is, in fact, transformed into work. Similarly, when the poles of the dynamo are properly connected and electricity is being produced, and only then, motion is used up. This is shown by the effort required to turn the armature under these circumstances, and the ease with which it is turned when the circuit is open. So, with a conductor like the filament in the lamp, unless it offers resistance to the current and de- stroys a sufficient amount of electrical energy, it gives out neither light nor heat. Finally, motion gives no heat unless the brake is set, and effort is then demanded to maintain the motion. These experiences lead us to believe that we have here a set of things which are fundamentally of the same kind, for each form can be made from any of the others. We have, therefore, invented the conception of a single thing, of which heat, light, electricity, and FIG. 54. 152 TEXT BOOK OF CHEMISTRY motion are forms, and to it we give the name energy: energy is work and every other thing which can arise from work and be con- verted into work (Ostwald) . Closer study shows that equal amounts of electrical or mechani- cal energy always produce equal amounts of heat. No loss is ever observed in the transformations of energy, any more than in the transformations of matter. Hence, J. R. Mayer (1842), Colding (1843), and Helmholtz (1847) were led independently to the conclusion that in a limited system no gain or loss of energy is ever observed. This brief statement of the results of many ex- periments is called the law of the conservation of energy. Application of the Conception of Energy in Chemistry. At first sight it looks as if the statement that energy is conserved is not applicable in chemistry. Heat and electricity, for example, seem to be produced and consumed, in connection with changes in composition, in a mysterious manner. We trace light in an in- candescent lamp back to the electricity, and this in turn to the mechanical energy, and this again to the heat in the engine. But what form of energy gave the heat developed by the combustion of the coal under the boiler, or by the union of iron and sulphur in our first experiment? Since we do not perceive any electricity, light, heat, or motion, in the original materials, and yet wish to create an harmonious system, we are bound to conceive of the iron and the sulphur, and the coal and the air, as containing an- other form of energy, which we call internal energy. Similarly, when heat is used up in decomposing mercuric oxide, or light in decomposing silver chloride, we regard the energy as passing into, and being stored in the products of decomposition in the form of internal energy. These conclusions compel us, for the sake of consistency, to think of all our materials as repositories of energy as well as of matter, each of these two constituents being equally real and equally important. A piece of the substance known as "iron" ENERGY AND CHEMICAL CHANGE 153 must thus be held to contain so much iron matter and so much internal energy. So ferrous sulphide contains sulphur matter, iron matter, and internal energy. Thus, by a substance we mean a distinct species of matter, simple or compound, with its appro- priate proportion of internal energy. In the course of this discussion it has become clear that it is characteristic of chemical phenomena that, besides a change in the nature of the matter, there is always an alteration in the amount of internal energy in the system. This alteration involves the produc- tion of internal energy from, or the transformation of internal energy into some other form of energy. Energy and Chemical Activity. Other things being equal, actions in which there is a relatively large loss of internal energy proceed rapidly; that is to say, in them a large proportion of the material is changed in the unit of time. Those in which less en- ergy is transformed proceed more slowly. The speed of the chemi- cal change, and the quantity of energy available because of it, are closely related. Now, we are accustomed to speak of materials which, like iron and sulphur, interact rapidly and with liberation of much energy as "chemically active." Thus, relative chemical activity may be estimated, (1) by observing the speed of a change (see below), or, in many cases (2) by measuring the heat developed (see Thermo-chemistry, below), or (3) by ascertaining the electro- motive force of the current, when the materials are arranged in the form of a battery-cell. It is evident that the chemical activity of a given substance will not be the same towards all others. Thus, iron unites much more vigorously with chlorine than with sulphur and, with identical amounts of iron, more heat is liberated in the former case than in the latter. With silver, sodium, and many other substances, iron does not unite at all. One of the tasks of the chemist is to make such comparisons as this. He calls the results, the specific chemi- cal properties of the substances in question. 154 TEXT BOOK OF CHEMISTRY The "Cause" of Chemical Activity or Affinity. The reader will undoubtedly be inclined to inquire whether we can assign any cause for the tendency which substances have to undergo chemical change. Why do iron and sulphur unite to form ferrous sulphide, while other pairs of elements taken at random will fre- quently be found to have no effect upon one another under any circumstances? The answer is that we do not know. Questions like this have to go without answer in all sciences. What is the cause of gravitation? We know the facts which are associated with the word the fact that bodies fall towards the earth, for example but, why they fall we are unable to say. So, with chemical change, we can state all the facts we know about it, but even then we cannot say why it takes place. The word cause was employed in the heading of this section, and it will be observed that no cause was found. This is the invariable rule in physical or chemical phenomena. We know of no causes, in the sense in which the word is commonly employed. The word cause has only one definite use in science. When we find that thorough incorporation of the three materials is needed to secure good gunpowder, we say that the intimate mixing is a cause of its being highly explosive. By this we simply mean that inti- mate mixture is a necessary antecedent of the result. A cause is a condition or occurrence which always precedes another condition or occurrence. The Speed of Chemical Actions: a Means of Measuring Activity. One means of measuring the relative chemical activi- ties of several substances is to observe the speed with which they undergo the same chemical change. Thus we may compare the activities of the various metals by allowing them separately to interact with hydrochloric acid and collecting and measuring the hydrogen liberated per minute by each. It will be seen, even in the roughest experiment, that magnesium is thus much more active than zinc. The comparison must be made with such precautions, ENERGY AND CHEMICAL CHANGE 155 however, as will make it certain that the conditions under which the several metals act are all alike. Thus, in spite of the heat evolved by the action, means must be used, by suitable cooling, to keep the temperature at some fixed point during the experiment, for all actions become more rapid when the temperature rises (p. 14). Again, the pieces of the various metals must be arranged so that equal surfaces are exposed to the acid in each case. It is found that the order in which this comparison places the metals is much the same as that in which they are placed by a study of other similar actions. A single table, showing the order of activ- ity (p. 54), suffices, therefore, for all purposes. Thermochemistry. Chemical changes in which heat is liberated are called exothermal. Those in which heat is continu- ously absorbed (pp. 15, 30) are called endothermal changes. Since the activities, or affinities of two substances (say, two metals) may often be measured by observing the amounts of heat liberated when each combines with a third substance (say, oxygen), it will be instructive now to consider some of the elementary facts of thermochemistry. The chemical interactions to be studied thermally are arranged so that they may be carried out in a small vessel which can be placed inside another containing water. The whole apparatus is called a calorimeter (Gk., heat-measurer). The heat developed raises the temperature of this water. Where gases like oxygen are concerned, a closed bulb of platinum forms the inner vessel. The quantity of heat capable of raising one gram of water one degree in temperature at 15 Centigrade is called a calorie. So that 250 grams of water raised 1 would represent 250 calories, and 20 grams of water raised 5 would represent 100 calories. Thermochemical Equations. While in physics the unit of quantity is the gram, in chemistry the unit which we select is naturally a gram-atomic weight or a gram-molecular weight of the 156 TEXT BOOK OF CHEMISTRY substance. Thus, the heat of combustion of carbon means the heat produced by combining twelve grams of carbon with thirty- two grams of oxygen, and is sufficient to raise nearly 100,000 grams of water one degree. This is expressed as follows: C -{- 2 -> CO 2 + 96,820 cal. In other words, the combustion of less than half an ounce of carbon will raise over two pounds (one kilogram) of water from to the boiling-point. When the action is one which absorbs heat, this fact is indicated by the negative sign preceding the number of calories. Thus, the dissociation of 36 g. of water vapor into hydrogen and oxygen absorbs 28,800 cal. per gram of hydrogen liberated: 2H 2 O -* 2H 2 + O 2 - 115,200 cal. If the action is reversible, as this one is, the heat absorbed when it proceeds in one direction is equal to that liberated when it goes in the other direction: 2H 2 + O 2 -> 2H 2 + 115,200 cal. An action which absorbs heat can take place only when heat or some other form of energy is furnished. Thus, the electrolysis of aqueous hydrochloric acid (p. 55) consumes electrical energy, which is equivalent in amount to the heat given out when hydrogen and chlorine unite to form hydrogen chloride, plus the heat liber- ated when the latter dissolves: H 2 + C1 2 + Aq. -> 2HC1, Aq. + 78,800 cal. Answers to Possible Questions. It is always found that the same quantities of any given chemical substances, undergoing the same chemical change under the same conditions, produce or absorb, according as the action is exothermal or endothermal, amounts of heat which are equal. The rate at which a given chemical action is allowed to take place has no influence on the total amount of heat consumed or produced. ENERGY AND CHEMICAL CHANGE 157 Tt may not at first sight appear obvious that rusting evolves heat, but a delicate thermometer will show that a heap of rusting nails is somewhat higher in temperature than surrounding bodies. Poor conductors, like oily rags and ill-dried hay, show a tendency to spontaneous combustion owing to accumulation of the slowly de- veloping heat of oxidation (p. 41). The warmth of our own bodies is due to the same cause. It should be noted that production or absorption of heat is not, in itself, an evidence of chemical action. Physical changes are all likewise accompanied by the same phenomena. Thus, the evapo- ration of water absorbs heat, and condensation of a vapor and the crystallization of a supercooled liquid liberate heat. The heat of solution (cf. pp. 112, 116) is the heat liberated (or absorbed) on dissolving one mole of the substance in a large amount of water. A part of the water always undergoes chemical change (p. 117). The solute also frequently combines with a part of the water, or is ionized (q.v.), and the change in volume of the mixture (p. 113), as a physical phenomenon, would alone entail a heat-change. Hence this heat effect is partly chemical and partly physical in origin. Exercises. 1. Which form of energy is delivered as such, and paid for as such, in most cities ? 2. How many calories are required to raise 500 g. of a substance of specific heat 0.5 from 15 to 37? 3. The combustion of 1 g. of sulphur to sulphur dioxide develops 2220 calories. What is the heat of combustion of sulphur ? Write the thermochemical equation. CHAPTER XIII SODIUM AND SODIUM HYDROXIDE IN our study of common salt we have taken up one of its con- stituents, namely, chlorine, and its commonest derivative hydro- chloric acid. The latter is a good example of an acid. We now turn to the other constituent, sodium, and one of its familiar compounds, namely sodium hydroxide (NaOH). The latter is an example of a different kind of substances, called bases. Salt, hydrochloric acid, and sodium hydroxide are examples of the three largest and most important classes of substances known to inorganic chemistry. Preparation of the Metal Sodium. Sodium cannot be made by displacement (like hydrogen), because it is close to the top of the activity list (p. 54) of the metals, and no more active, and at the same time easily ob- tained, metal is available to displace it. It was first prepared by Davy (1807) by electrolyzing melted sodium hydroxide, and is still manufactured in this way. The aqueous solution of a sodium compound, such as sodium chloride, cannot be used, because, as we have seen (p. 133), hydrogen from the water is liberated in place of the metal. The dry sodium hydroxide is melted (about 318) in an iron vessel (Fig. 55), which is connected with the positive wire from the dynamo, and the oxygen is liberated on the iron (anode). The negative wire is connected with rods of carbon which stand up through the bottom of the vessel, and here (cathode) the sodium and the hydrogen are set free. Being lighter than the fused com- pound, the hydrogen rises in bubbles, and the sodium in melted 158 SODIUM AND SODIUM HYDROXIDE 159 globules, to the surface. Here they collect under an iron cylinder. The latter is made of wire-gauze at the lower part, to permit cir- culation of the liquid, but prevent the escape of the globules of sodium. It is closed at the top, to prevent the heated sodium from burning, as it would do if air could reach it. The melted sodium is ladeled into cylindrical moulds, and the sticks of the metal are preserved in air-tight tin boxes. Physical Properties of Sodium. Sodium is a silver-like metal, of specific gravity 0.97. It is soft and can be cut with a knife. It melts at 95.6 and boils at 742. The gram-molecular volume of sodium vapor weighs 23 g., the same as the atomic weight, so that the molecular formula is Na. Chemical Properties. Sodium burns in chlorine, giving sodium chloride NaCl. It burns also in oxygen (or air) to form sodium peroxide NaA: It acts violently on water, as we have seen (p. 50), displacing hydrogen and forming sodium hydroxide. Skeleton: Na + H 2 O -> NaOH + H 2 . Balanced: 2Na + 2H 2 O -> 2NaOH + H 2 . For this reason it tarnishes quickly in moist air. In the labora- tory small amounts are kept under kerosene, which contains no compounds of oxygen. As shown by its place in the activity list (Appendix V), the metal tends to form very stable compounds, and displaces the other metals (except potassium) from combination. Uses. Sodium is used in the manufacture of many complex organic compounds. By contact action, it converts isoprene (C 5 H 8 ) into caoutchouc (C 5 H 8 ) 4 or raw rubber. This is a method of making rubber artificially, but it has not yet been used on a commercial scale. Sodium peroxide ("oxone")> made by burning sodium, is used as a source of oxygen (p. 31). 160 TEXT BOOK OF CHEMISTRY Preparation of Sodium Hydroxide. As we have seen, this compound is formed by the action of water on sodium, and may be obtained by evaporating the solution to dry ness. But man- ufacturing it from an expensive substance like sodium is out of the question. Much sodium hydroxide is made by boiling an aqueous solution of sodium carbonate Na 2 C03 with slaked lime (calcium hydroxide, Ca(OH) 2 ): Ca(OH) 2 + Na^COs <=* 2NaOH + CaCO 3 . The calcium carbonate, which is chemically the same substance as lime-stone and chalk, is insoluble. All double decompositions are reversible, but here the precipitation of the calcium carbonate prevents it from acting upon the sodium hydroxide and reversing the action. After the precipitate has settled, the solution of sodium hydroxide is drawn off and evaporated to dryness. Electrolytic Process. Sodium hydroxide is also manufac- tured by electrolysis of sodium chloride solution (pp. 133-4), the other product, chlorine, being of great commer- cial value also. In the Castner-Kellner appara- tus (Fig. 56) the sodium chloride solution occupies the end compartments. The chlorine is liberated FIG. 56. upon the carbon electrodes in these compartments and rises to the surface. It is drawn off, dried, and compressed to liquid form in iron cylinders, or is made directly into bleaching compounds (p. 307). The sodium passes to the layer of mercury upon the floor of the box and dissolves in it. The box is caused to rock by the cam (C), so that the mercury (alone) flows under the partitions. In the center compartment the sodium leaves the mercury and proceeds towards the negative electrode which hangs SODIUM AND SODIUM HYDROXIDE 161 here. The hydrogen of the water, being more easily liberated (as its place in the activity list shows), is set free and sodium hydroxide remains: Skeleton: Na + H 2 O - NaOH + H 2 Balanced: 2Na + 2H 2 O - 2NaOH + H 2 . The solution is evaporated as usual. More salt can be added in the end compartments from time to time, so that the process is continuous. Physical Properties of Sodium Hydroxide. The substance is a white crystalline solid. Generally it shows the form of the iron drums, into which it is run when melted, or of the sticks into which it is cast. It is exceedingly soluble in water. Its solution gives to objects the smooth, soapy feeling which is characteristic of alkalies. The solution is sometimes called soda-lye, and the solid, caustic soda. Chemical Properties. Sodium hydroxide is exceedingly stable, being melted, but not decomposed, by heating. The aqueous solution has an acrid taste, like soap or borax a property common to all alkalies. The solution acts upon azolitmin (contained in litmus ex- tract or litmus paper), after it has been altered by a trace of acid into red litmus, and changes the color back to deep violet- blue. This property is also shown by all alkalies. Sodium hydroxide in solution enters into double decomposition with many substances, and, when one of the products is insoluble, the action is seen to be instantaneous. For example, with cupric chloride, it gives a precipitate of cupric hydroxide. Skeleton: NaOH + CuCl 2 -> Cu(OH) 2 j + NaCl. Balanced: 2NaOH + CuCl 2 -> Cu(OH) 2 j + 2NaCl. As this equation shows, sodium hydroxide behaves in such actions as if composed of (Na) and (OH). These actions, therefore, 162 TEXT BOOK OF CHEMISTRY afford a ready means of making hydroxides of other metals. This property, again, is common to all alkalies. Solutions of the alkalies also act upon animal matter, e.g. wool (p. 1), especially when hot, converting it largely into soluble substances. For this reason they are called caustic alkalies. They likewise act slowly upon the components of glass. For this reason a precipitate is often visible in the caustic soda reagent bottle, and the inner surface of the glass if always etched. Alkalies and Bases. It will be seen that the chemical proper- ties of sodium hydroxide may be summed up by saying that it is an alkali. The alkalies, however, are simply the more soluble mem- bers of a much larger class of substances called bases. The less soluble bases, of which cupric hydroxide is an extreme ex- ample, do not show, distinctly, all the properties exhibited by alkalies. Thus, those which are least soluble have, naturally, no taste, do not visibly affect litmus, and are not soapy to the touch or corrosive towards glass. But they all show the tendency to double decomposition, in which the group (OH) is transferred, as it was from NaOH to Cu(OH) 2 in the foregoing example. Uses.* Sodium hydroxide is used in immense quantities along with fats, in the manufacture of soap. Some bleaching liquids are made by saturating it with chlorine. It is employed also in making many other sodium compounds which are used in the arts. Exercises. 1. In the electrolysis of sodium hydroxide, why is the metal not liberated as a solid or as a vapor, but as a liquid ? 2. Make molecular equations for the burning of sodium in chlorine and in oxygen. 3. Make equations showing the interactions of solutions of aluminium chloride A1C1 3 and of cupric sulphate CuS0 4 with sodium hydroxide (p. 161). Name the products. CHAPTER XIV ACIDS, BASES, AND SALTS As we saw, at the opening of the preceding chapter, acids, bases, and substances with certain chemical properties similar to those of common salt form the vast majority of the compounds met with in inorganic chemistry. On account of their resemblance to common salt, the substances of the third class are called salts. We shall first recall the names and formula of some of the examples of these classes that we have already had occasion to mention. Some Familiar Acids. The following are the acids we have mentioned : Hydrochloric acid HC1 Nitric acid .HNO 3 Sulphuric acid H 2 SO 4 Hypochlorous acid. . . . HC1O Phosphoric acid H 3 PO 4 Acetic acid ,SGO 2 CH, Hydrofluoric acid HF Acids are sometimes divided into organic acids, like acetic acid, oxacids, like sulphuric acid, and hydracids, like hydrochloric and hydrofluoric acids, which contain no oxygen. Familiar Bases. The bases mentioned, with one or two addi- tions (in italics), are: Sodium hydroxide NaOH Calcium hydroxide Ca(OH) 2 Potassium hydroxide KOH Cupric hydroxide Cu(OH)2 Ammonium hydroxide. . . .NH 4 OH Zinc hydroxide Zn(OH)j The name alkalies (p. 162) is applied to all the three soluble ones in the first column, and often to calcium hydroxide, although it is much less soluble. 163 164 TEXT BOOK OF CHEMISTRY Familiar Salts. Some of those already mentioned are: Sodium chloride NaCl Potassium chlorate .... KClOa Sodium sulphate Na 2 SO 4 Zinc sulphide ZnS Potassium nitrate KN0 3 Sodium peroxide Na 2 O 2 Ammonium chloride NH 4 C1 Sodium carbonate Na 2 COa Cupric sulphate CuSO 4 Lead sulphide PbS Cupric chloride CuCl 2 Mica Zinc chloride ZnCl 2 Felspar Zinc sulphate ZnSO 4 Glass Radicals. Upon comparing the formulae in the lists, we notice that: 1. Every acid contains hydrogen, combined with one or more other atoms, constituting a group called a radical. 2. The very same radicals, such as S0 4 , appear also in a number of salts. 3. Every base contains OH (the radical hydroxyl), and usually one atom, but occasionally more than one (as in NH 4 ) in the other radical. 4. The same radicals, combined in bases with OH, appear also in numerous salts. 5. Radicals may be simple, like H, Na, and Cl, or compound, like SO 4 and NH 4 . Since, in electrolysis, the hydrogen of an acid is attracted to, and is liberated at the negative electrode, it is called a positive radical. The rest of the molecule, as Cl or SO 4 , is the negative radical. Similarly, every base contains OH as the negative radical, combined with a positive radical, as Na or NH 4 . Again, every salt contains a positive radical, other than H, combined with a negative radical, other than OH. The name of each salt indicates the radicals of which it is composed*. As regards the .salts in the second column, they show positive radicals, like K, Zn, and Na found in the list of bases. The negative radicals, C10 3 , S, 2 , and C0 3 are not in the list of acids we have met with. But there are well-known acids correspond- ACIDS, BASES, AND SALTS 165 ing to them, namely, chloric acid HC10 3 , hydrogen sulphide H 2 S, hydrogen peroxide H 2 O2, and carbonic acid H 2 CO 3 . In general, then, ah 1 positive radicals combine with OH to give bases, all negative radicals combine with H to give acids. In general, also, each positive radical will combine with any negative radical to give a salt. In a few exceptional cases only, the compound can not be formed, presumably because it is un- stable under ordinary conditions. Reaction Formulae. The formulae of acids, bases, and salts are written in a uniform manner to show the behavior of the substances represented. Thus, the radical written first is usually the positive one which, in electrolysis, is attracted by the nega- tive wire, and the negative radical follows. Again, in a compound like calcium hydroxide, the formula CaO 2 H 2 would conceal the existence of the hydroxyl group. So the radicals are written in brackets, with the coefficient outside Ca(OH) 2 . Thus we write also Cu(N0 3 ) 2 , and not CuN 2 O 6 , and (NH 4 ) 2 SO 4 , not N 2 H 8 SO 4 . These formulae are all reaction formulae. That is, they indicate, not simply the composition, but also the parts into which the compounds decompose, and from which they are formed in double decompositions (p. 161, and p. 166, below). Other Classes of Substances. The classification into the three groups, acids, bases, and salts, clears the ground wonderfully, for the substances we shall meet with, that are not included in one or other of these groups, are very few indeed. Of the substances not belonging to those classes, the following may be mentioned. 1. The elementary substances, of which there are about eighty. Of these we have mentioned a dozen. 2. The oxides, of which we have met some fifteen, including stannic oxide SnO 2 , mercuric oxide HgO, and cupric oxide CuO. Theoretically, the oxides might be classed as salts, with the negative radical 0. Water H 2 is the corresponding acid. But 166 TEXT BOOK OF CHEMISTRY there are some objections to this view, and it is not at present adopted by chemists. 3. The organic compounds, of which carbon disulphide 082, formaldehyde CH 2 0, alcohol C2H 6 O, turpentine CioHi 6 and others have been mentioned. There are many organic acids, bases, and salts. But the examples just named, as well as a vast number of other organic compounds, do not belong to any one of these three classes. 4. The molecular compounds, to which the hydrates, like CuS0 4 , 5H 2 O belong. Properties Common to Acids, Bases and Salts. There are four of these properties, three of which have come up repeatedly. These three form the subject of this section. (For the fourth property, see p. 172.) All the four are chem- ical properties, that is, they involve some form of chemical changes. 1. Double Decomposition. In solution, an acid or a base will interact with any salt. Acids also interact with bases and salts with salts. In all cases the action is a double decomposition (defined in p. 126). For example: Salt and acid: AgN0 3 +HC1 -> AgCl J, + HN0 3 . Salt and base: Na*CO 3 + Cu(OH) 2 -> CuCO 3 j + 2NaOH. Acid and base: 2HC1 +Ca(OH) 2 -+ CaCl 2 + H 2 0. Salt and salt: 2AgN0 3 +CaCl 2 - 2AgCl J, + Ca(N0 3 ) 2 . In each case a simple exchange of radicals occurs. When one product is insoluble, the action is instantaneous and complete. In the third equation, where an acid and base interact, one of the products is necessarily water, obtained by the union of the H and OH. This case is called neutralization, for the special properties of the acid and base, such as the tastes and the actions upon litmus, disappear. Many of these actions are used as tests for the radicals in un- ACIDS, BASES, AND SALTS 167 known substances, because the appearances and colors of the precipitates differ more or less distinctly (p. 126). These actions also furnish us with a general method of making every substance of three classes, namely, by choosing two other suitable substances containing the required radicals. A radical is an atom, or group of atoms, which behaves as a unit in double decomposition. 2. Displacement. The simple radicals can be displaced, and the element liberated. Thus, we prepared hydrogen by adding metals to acids, and we absorbed chlorine by contact with potas- sium iodide solution. Zn + 2HC1 -> H 2 + ZnCl 2 2K1 + Cl - 2KC1 + I 2 . This does not apply to compound radicals. 3. Electrolytes. Solutions of acids, bases, and salts in water are all conductors of electricity. In all cases the solute is decom- posed by the current. The separation is always into the two radicals, which are carried to opposite electrodes. Thus, all acids give hydrogen at the negative pole, the other radical passing to the positive pole. Of all the properties they have in common, this one of being electrolytes is perhaps the most remarkable. It appears more surprising when we consider that no substances, other than acids, bases, and salts, undergo electrolysis in aqueous solution. This is an exclusive property of these classes of bodies. It should be noted also that the water evidently has a good deal to do with the process. Alcoholic solutions behave in the same way. But solutions in benzene and carbon disulphide and many other solvents are non-conductors. The acids, bases and salts, when pure and free from water, are also non-conductors. The Conductivity of Electrolytes. The conducting power of solutions can be examined, roughly, by the apparatus in Fig. 57. 168 TEXT BOOK OF CHEMISTRY FIG. 57. The platinum electrodes are connected with a direct-current cir- cuit. The lamp, which is on one of the wires, serves, by its resistance, to cut down the current. Its glowing also shows when the liquid is a conductor, and by varying brightness indicates roughly the conducting power of the solution. A solution of sugar in water shows no conductance, and the lamp remains dark. Solu- tions of acids, bases, and salts in water enable the lamp to glow. We quickly find that differ- ent substances, when they conduct, do so in different degrees. Hydrochloric and nitric acids conduct well. So do sodium and potassium hydroxides. Salts are practically all good conductors. But many acids, like acetic acid, conduct poorly. The same is true of some bases, like am- monium hydroxide. Of course, insoluble substances give poorly conducting solu- tions. But in estimating the conductivity for chemical pur- poses, we take into consideration the amount dissolved. Thus silver chloride, being a salt, is an excellent conductor, when we allow for the extreme diluteness of the solution. The unit is the conducting power of one molecular weight of the substance, placed in solution between the electrodes. Reasoning still further about these phenomena we see that, if we start with a fixed amount of a given substance, the conductivities at different stages of the dilution must be proportional to the numbers of ions, and the maximum conductivity attainable by great dilu- tion must represent the effect when the whole material has become ionic. Thus, if the conductivity at the maximum is represented, say, by 5, then at the dilution where the conductivity is 2, the proportion of the whole which is ionized is 2/5. When the con- ductivity becomes 4, 4/5 of the molecules are dissociated and the ACIDS, BASES, AND SALTS 169 degree of ionization is 0.8. When the conductivity becomes 5, 5/5, or all, of the molecules are dissociated. For example, in hydrochloric acid, if we take the normal solution (p. 115) containing 36.5 g. of acid per liter as the unit of concentration, the fractions ionized at various concentrations are as follows: 1(X/V, 0.17; N, 0.78; N/W, 0.91; AT/100, 0.96. Thus, measurements of con- ductivity enable us to study the ionic decomposition of all ionogens, and to state accurately the fraction ionized, at each concentration, in solutions of every ionogen. This information is obviously most valuable, for it places us in a position to know the exact constitu- tion of every solution we use in the laboratory. In the following section the data on which such knowledge can be based is given. In the next chapter the mode of applying the data is explained. Constitution of Solutions of Ionogens: Fractions Ionized. The dilute acids used in the laboratory are generally of six times normal (QN) concentration. But, often, we add only a drop or two to a large bulk of liquid, so that the acids are commonly very dilute as actually employed. The solutions of salts are of different strengths, but the great majority are of normal (N), or even smaller concentrations. In practice they, also, are still further consider- ably diluted before use. If, therefore, we give the fractions ionized (total molecules of ionogen = 1) in decinormal solutions (except where otherwise specified), the reader will be able to estimate roughly the proportion of each kind of ions in any application of the reagent. In the case of acids containing more than one dis- placeable hydrogen unit, the kind of ionization on which the figure is based is indicated by a period. Thus H.HCOs means that the whole of the ionization is assumed to be into H + and HG0 3 ~. 170 TEXT BOOK OF CHEMISTRY FRACTION IONIZED IN 0.1 AT SOLUTIONS AT 18 Acros Nitric acid 0.92 Nitric acid (cone., 62%) . . 09 Hydrochloric acid . . .0.92 Hydrochloric acid (cone., 35%) 0.13 Sulphuric acid, H.H.SO 4 . .0.61 Sulphuric acid (cone., 95%) 0.01 Hydrofluoric acid . . . .0.15 Oxalic acid, H.HC 2 O 4 . . .0.50 Tartaric acid, H.HT . . .0.08 Acetic acid (AT) 0.004 Acetic acid . 0.013 Potassium hydroxide . . . . 91 Sodium hydroxide . . . . . 91 Barium hydroxide . . . . . 77 Lithium hydroxide (N) . . . 63 Tetramethylammonium hy- droxide (AT/16) .... 0.96 Potassium chloride . . Potassium nitrate . . Potassium acetate . . Potassium sulphate Potassium carbonate Potassium chlorate Ammonium chloride . Sodium chloride (AT) . Sodium chloride (AT/2) Sodium chloride . . . Sodium nitrate. . . . Sodium acetate . . . Sodium sulphate . . . . 0.86 . 0.83 . 0.83 . 0.72 (0-71) . 0.83 . 0.85 . 0.66 . 0.74 . 0.84 . 0.83 . 0.79 . 0.70 Carbonic acid, H.HC0 3 . - 0.0017 Carbonic acid (AT/25) . . . 0.0021 Hydrogen sulphide, H.HS . 0.0007 Boric acid, H.H 2 BO 3 .... 0.0001 Hydrocyanic acid 0.0001 Permanganic acid (A r /2) . . . 93 Hydriodic acid (N/2) . . .0.90 Hydrobromic acid (N/2) . . 90 Perchloric acid (AT/2) . . .0.88 Chloric acid (N/2) 0.88 Phosphoric acid, H.H 2 P0 4 . .0.27 Water . . 0.0 6 1 BASES Ammonium hydroxide . . 0.013 Strontium hydroxide (AT/64) 0.93 Barium hydroxide (AT/64) .0.92 Calcium hydroxide (AT/64) . . 90 Silver hydroxide (AT/1783) .0.39 Water . . 0.0 6 1 SALTS Sodium bicarbonate, Na.HC0 3 0.78 Sodium phosphate, Na 2 .HP0 4 0.73 Sodium tartrate . 69 Barium chloride . 77 Calcium sulphate (N/100) .0.64 Cupric sulphate . 39 Silver nitrate 0.81 Zinc sulphate . 40 Zinc chloride 0.73 Mercuric chloride . . . ( < . 01) Mercuric cyanide Minute In additon to their use in showing the nature of the reagents employed in the laboratory (p. 168), these numbers show also to what extent any pair of ionic substances will unite when mixed (see ACIDS, BASES, AND SALTS 171 pp. 303-305), and they likewise indicate the chemical activity of the ionogens when in solution (see next section). Relation of lonization to Chemical Activity. These tables may be used for reference. The import of the following general statements, drawn from the tables, should be memorized: 1. Salts, with the exception of those of mercury, are all well ionized. In actions involving their ions, salts are therefore all of the same order of activity, for a dilute solution of every salt con- tains a large amount of the ionic components. 2. Acids show the most extreme differences in their degrees of ionization. That is to say, their solutions must contain very differ- ent concentrations of hydrogen-ion. Since then* activity as acids depends on this substance (p. 170), and the activity of a substance is proportional to its concentration (p. 232), it follows that acids will show very great differences in apparent chemical activity. At this point, therefore, we emerge from semi-physical discussion of the subject and reach something of definite, practical application in chemical work. The data show that acids may be divided roughly into four classes with different degrees of acidic activity: (a) The ionization in decinormal solution exceeds 70 per cent; e.g., nitric acid and hydrochloric acid. These are the acids which are chemically most active, for their solutions contain a relatively high concentration of hydrogen-ion. (b) The ionization is between 70 and 10 per cent; e.g., sulphuric acid and phosphoric acid. These acids are noticeably less active, for their solutions contain a lower concentration of hydrogen-ion. (c) The ionization is between 10 and 1 per cent; e.g., acetic acid. These are the weaker acids, for their solutions contain a very small concentration of hydrogen-ion. (d) The ionization is less than 1 per cent; e.g., carbonic and boric acids. These are the feeble acids, for their solutions contain only a minute concentration of hydrogen-ion. 172 TEXT BOOK OF CHEMISTRY 3. The bases show two classes: (a) lonization high; e.g., potassium hydroxide. These bases are active, for their solutions contain a high concentration of hydroxide-ion. (b) lonization less than 2 per cent; e.g., ammonium hydroxide. These bases are weak on account of the low concentration of hydroxide-ion. 4. Water is less ionized than any other substance in the list. It shows therefore, as we already know, usually little or no interaction with acids, bases, or salts, and hence is valuable as a solvent for these substances. Its ions are H + and OH~, and it is thus as much an acid as a base. Boiling- and Freezing-Points of Solutions of Acids, Bases, and Salts. We have seen (p. 112) that equal numbers of mole- cules dissolved in equal volumes of water raise the boiling-point to the same extent, and lower the freezing-point by a constant amount. Thus, one molecular weight of sugar (342 g.) or of glycerine (92 g.) dissolved in 1000 c.c. of water will raise the boiling-point from 100 to 100.52, and will lower the freezing- point from to 1.86. But this is uniformly true only of sub- stances which are not acids, bases, or salts. Molecular weights of substances of these three classes, in 1000 c.c. of water, raise the boiling-point more than 0.52 degrees and lower the freezing-point by more than 1.86 degrees. We say they give abnormal elevations of the boiling-point and abnormal lowerings of the freezing-point. Thus, a solution of 58.46 g. of sodium chloride in 1000 c.c. of water, boils at 100.87, and freezes at -3.65. The elevation in the boiling-point of the water is 0.87 degrees instead of 0.52 degrees. The effect is nearly twice as great as the normal one. The only conclusion we can draw is that nearly twice the usual number of molecules is present. In other words, most of the molecules of the sodium chloride (but not all) have been broken into two parts by the water, These sub-molecules are called ions. ACIDS, BASES, AND SALTS 173 The decomposition into ions is a chemical change. It is the fourth property (p. 167) of acids, bases, and salts. Further, we find that the abnormal effect varies in amountj yet not in an erratic manner, as the two following facts show: Jons and Conductivity. The electrolytes which are the best conductors of electricity, such as hydrochloric acid, sodium hydroxide, and salts, produce the greatest abnormal effects upon the boiling- and freezing-points. On the other hand, poor con- ductors, like solutions of acetic acid and ammonium hydroxide, give almost normal elevations and depressions. Careful measure- ments show a close agreement in respect to conductivity, on the one hand, and elevation or depression on the other, throughout the whole list of hundreds of substances. This cannot be a chance coincidence, and demands explanation. It shows that conductivity depends upon the parts of molecules we have called ions, and is pro- portional to their number. Jons and Radicals. The other fact is that, while substances like sodium chloride (NaCl) and nitric acid (HNOs), which con- tain only two radicals, give effects approaching double the normal ones, they never give more than twice the normal effect. But sub- stances like sodium sulphate Na2SO4 and calcium nitrate Ca(NOs)2, which contain three radicals, give effects approaching three times the normal, but not more. And substances containing four radicals, like ferric chloride FeCl 3 , give effects approaching four times the normal. Evidently, solution in water tends to separate most of the radicals from one another, and to break each molecule into as many separate particles as there are radicals in the sut> stance. This indicates that the radicals separate to form the ions. Apparently, therefore, the true basis of the four properties which are common to acids, bases, and salts is that they all separate into ions when dissolved in water. Hence, to save repeating the three words, we may name these substances, collectively, ionogens, 174 TEXT BOOK OF CHEMISTRY from the Greek, meaning ion-producers. The decomposition into ions is called ionization. lonization a Reversible Action. The depression in the freezing-point produced by sodium chloride was 3.65 degrees, instead of 1.86 degrees. A depression of double the normal value would have been 2 X 1.86 = 3.72. The decomposition was there- fore incomplete. A decomposition of 67 per cent of the mole- cules would give the observed depression of 3.65. We find, however, that the addition of more water increases the percentage of ionized material. This is shown by measuring boiling-points, freezing-points, or conductivities. On the other hand, removal of water, by evaporation, causes some of the ions to unite. Finally, when all the water is gone, the solid that remains is, of course, composed wholly of molecules. lonization is therefore a reversible chemical action. NaCl^(Na) + (Cl). Ions Bear Electrical Charges. We have yet to explain twc facts, namely, (1) how it is that ionogens in aqueous solution are decomposed by the current in electrolysis, and (2) how they have a power of conducting electricity proportional to the extent ol their ionization (p. 169). The explanation, first offered by the Swedish chemist Arrhenius (1887), is startling, but satisfactory, It is, simply, that the ions bear electrical charges. In this respect they differ from ordinary molecules or atoms. The constituents of the acid, base, or salt are liberated because the negative electricity on the negative electrode attracts all the positive ions in the solution,' and repels all the negative ones. 4 At the same time, the negatively charged ions are attracted towards the positive electrode, and the positive ones repelled from it. The instant, therefore, that the circuit is completed all the ions in the solution begin to move, or migrate, as it is called * In all circumstances, unlike charges of electricity attract one another and like charges repel one another. ACIDS, BASES, AND SALTS 175 towards their proper electrodes (Fig. 58). Ions (Greek, going) derive their name from this fact. The rest is easily understood. When a positive ion, say, of hydrogen, reaches and touches the negative electrode, its positive charge of electricity is neutralized, and the result is an ordinary atom of hydrogen. The atoms of free hydrogen unite to give molecules (H 2 ) and these form bubbles of the gas. If the posi- tive ion is copper (the solution being one of cupric sulphate, let us say), then the copper atoms, when discharged, adhere to the electrode, and the latter becomes plated with copper. ^ ' ' 5 ' 8 Simultaneously with these operations, the negative ions are discharged at the positive elec- trode, and their material liberated. As to the conductivity of the solution, we must remember that pure water does not conduct the current, except to an extraordinarily small extent. The conducting power of solution is due, there- fore, to the ions. The more numerous the ions are, the greater will be the number of discharges taking place per second upon each electrode. So that the conducting power is naturally pro- portional to the number of ions, as we find it to be. It must be noted that the electricity does not flow, as it does in a copper wire. In the wire the material is at rest, and the electric- ity flows through it. In the electrolyte the particles of the sub- stance themselves move, and carry electricity with them. Ionic Equations. lonization is a kind of decomposition or rather dissociation, and is in every sense a chemical change. It is therefore expressed by equations . In these the charges upon the ions must be shown, as they are essential parts of the ionic substances. NaCl <= Na+ + Cl- FeCl 3 <=* Fe+++ + 3C1~ Ca(N0 3 ) 2 ^ Ca++ + 2NOr Na^ < 2Na+ + S0 4 ". 176 TEXT BOOK OF CHEMISTRY The number of positive charges must equal the number of negative charges. This is proved by several facts, and most simply by the fact that the solution of a substance like ferric chloride (FeCl 3 ) is electrically neutral, as a whole. Thus some ions carry one charge, like Na + and Cl~ and H + , others two charges, like Ca* 4 ' and SC>4 = , and so forth. An ion is an atom, or group of atoms, carrying an electric charge. The positive ions are called the cations, since they move to- ward the cathode. The negative ions are the anions, and move toward the anode. We can usually tell which is the positive radical in a formula because it generally consists of one atom of a metallic element (K + , Cu" 1 " 1 ", etc.) or of hydrogen. The negative radical may contain a metal, like the Mn in K(Mn04), but always along with a non-metal like oxygen. Possible Misunderstandings. Metallic sodium acts vio- lently upon water. But the theory of ionization assumes that sodium atoms, in the form of ions (Na + ), can exist in water. The heavy electrical charge makes sodium-ion an entirely different substance from sodium. When an ionogen is dissolved in water, a part dissolves to give molecules, and only a part gives ions. The solution of an ionogen is a mixture of water with three other (one molecular and two ionic) substances. The ions are the most active materials in the mixture. It may be asked why ions with opposite charges do not unite. The answer is that they do unite. But other molecules dissociate, so that the average number of ions and molecules remains the same. The charges are not derived from an electric current. They are upon the ions, as soon as the latter are formed, and are present at all times in all solutions of ionogens. The ions are formed by dissociation of the molecules, as quickly as the latter dissolve. ACIDS, BASES, AND SALTS 177 lojis and Double Decomposition. What we have just learned leads us to realize that there is, behind ordinary double decomposition, a more complex machinery than we had imagined. The ions are the active substances. Thus, in making sodium hydroxide, the equation Ca(OH) 2 + Na*C0 3 -* 2NaOH + CaCO 3 | has to be modified to show the ions: Ca++ + 2OH- + 2Na+ + C0 3 -> 2Na+ + 20H~ + CaC0 3 | When such solutions are mixed, the ions with opposite charges attract one another, so that some molecules of CaC0 3 are formed. This substance being insoluble, the molecules come out of solution, and more are formed, until the ions Ca++ and CO 3 ~ " are prac- tically all used up. This is the mechanism when precipitation occurs. When the mixture is filtered, and the filtrate is evaporated, the Na + and OH~ unite in increasing numbers, and finally the solid NaOH is obtained. In neutralization (p. 166), an acid and a base give water and a salt. Na+ + OH- + H+ + Cl- -> H 2 O + Na+ + Cl~ The H + and OH~ unite because, as we have seen, water is prac- tically a non-conductor, and its ions can not exist together, save in very small numbers. Thus a double decomposition may go to completion (without any precipitation) if the two interacting substances provide the ions of a substance which is very little ionized. The essential part of all neutralizations is contained in the equation: H+ + OH-->H 2 0. During subsequent evaporation, the ions of the salt (here Na + and Cl~) unite, until all the water, including that produced by the action, is gone and crystals of the salt remain. 178 TEXT BOOK OF CHEMISTRY . , Ions and Displacement. When a metal acts upon a dilute acid, and hydrogen is liberated, the ions are again concerned in the mechanism of the action: Zn + 2HC1 -> ZnCl 2 + H 2 f Zn + 2H+ + 2C1- -> Zn++ + 2C1~ + H 2 | From this it appears that each zinc atom simply takes the positive charges from two hydrogen ions of the acid. The discharged hydrogen then escapes. Thus all displacements are of the form: Zn + 2H+ - Zn++ + H 2 . Properties Peculiar to Acids. An acid has the four proper- ties possessed by all ionogens, with the additional feature that one radical is always H. Thus: 1. With bases and salts it gives double decompositions, using H as one radical. 2. Its hydrogen radical is displaced by certain metals. 3. Its solutions are electrolytes and hydrogen is liberated at the cathode: 2H+ + 2 - H 2 . 4. It is ionized in aqueous solution. If sufficiently soluble, it has, in addition, a sour taste and a special action on litmus (turning it red, a test first discovered by Boyle), both due to the hydrogen radical. Many other substances, such as sugar Ci 2 H 22 On, contain hy- drogen, but lack these properties, and therefore do not contain hydrogen as a radical, and are not acids. Briefly, an acid is a substance which contains hydrogen as a posi- tive radical, and, if so uble, has the sour taste and the action on litmus of this radical. An active or strong acid is one which is very largely ionized, giving much hydrogen-ion (like HC1), and therefore shows the properties of the radical H very markedly. A weak acid is one which is very little ionized, like acetic acid. ACIDS, BASES, AND SALTS 179 Properties Peculiar to Bases. The base has the four proper- ties of an ionogen, one of its radicals being hydroxyl OH: 1. In double decomposition with acids and salts it exchanges OH for another negative radical. 2. In electrolysis, the OH liberated, at the anode, gives water and oxygen 40H- + 4 -> 2H 2 O + 2 . 3. Its positive radical is displaced by other metals. 4. It is ionized in aqueous solution. If sufficiently soluble, it shows, when dissolved, the soapy taste and feeling, and the special action on litmus (turning red to blue), both due to the OH radical. Many other substances (such as sugar C^H^On) contain OH, but it is not the presence of these elements, but their presence as a radical that makes a substance a base. Briefly, a base is a substance which contains hydroxyl (OH) as a negative radical and, if soluble, has the taste and action on litmus of this radical. Strong bases, like sodium hydroxide, and weak ones like ammonium hydroxide, differ in the same way that strong and weak acids differ. That is, weak bases are little ionized, and show the properties of a base less markedly. Activity of Metals and Strength of Bases. It will be observed that the most active bases (alkalies) are hydroxides of the three metals (K, Na, Ca) which come first in the activity-list (p. 54). We may therefore judge the activity of a metallic ele- ment, at least roughly, either by trying its displacing power as a metal, or by measuring the activity (degree of ionization) of its hydroxide. Salts. As we have seen, salts are substances containing a posi- tive and a negative radical, other than hydrogen and hydroxyl. Some salts do indeed contain (H) or (OH) as a radical, but always in addition to two other radicals. Thus we obtained 180 TEXT BOOK OF CHEMISTRY sodium-hydrogen sulphate, NaHS0 4 , which gives H + , as well as Na + and S0 4 ~~. This is an acid salt. Basic calcium chloride Ca(OH)Cl is a basic salt. Exercises. 1. By what experiments should you determine which were the radicals in substances of the following composition: Cu(N0 3 ) 2 , CaC0 3 , NH 4 Br, NH 4 I, KC10 4 and KMn0 4 . There are always two ways, and usually three, of determining the radicals what are they (p. 177) ? 2. How should you determine whether a given substance were an acid, base, or salt (ionogen), or not? 3. When do the ions receive their charges of electricity? 4. Using the models given in p. 175, make the ionic equations representing the ionization of all the acids, bases, and salts, the formulae of which are given in pp. 178, 179. 5. Which are the anions and which the cations in the substances whose formulae are given on p. 170? 6. Using the models in pp. 178, 179, make equations for the liberation of copper and of chlorine by electrolysis. 7. Make an ionic equation (p. 178) for the displacement: (a) of hydrogen from dilute hydrochloric acid by magnesium; and (b) of copper from cupric sulphate solution by zinc. CHAPTER XV VALENCE THE differing number of charges on different ions has called our attention vaguely to a subject which must now be explored and set forth more clearly. Valence. The formulae of a number of common compounds, including some that we have met with, are as follows: NaCl ZnCl 2 SbClg SnCU NaBr ZnBr 2 SbBr 3 SnBr 4 Nal ZnI 2 SbI 3 We observe that one atomic -weight of sodium appears to unite with only one unit of another element, one unit weight of zinc with only two units of another element, a unit of antimony with not more than three, and a unit of tin with only four units. It seems that an atomic weight of each element has a fixed capacity for combining with not more than a certain number of atomic weights of other elements. Other compounds of only two elements have the formulae: HC1 H 2 NH 3 CEU Li 3 N Apparently Cl combines with only one atom of another element, O with two atoms, N with three atoms, C with four atoms. Also an atom of hydrogen combines with not more than one atom of another element, although it may take more than one atom of hydrogen to satisfy the atom of that other element (H 2 O, CHi, etc.). This limited combining capacity of each kind of atomic weight (or atom) is called its valence. 181 182 TEXT BOOK OF CHEMISTRY Marking the Valence. Until we are familiar with their values in each case, it may be well to mark the valences thus: Na 1 Zn 11 Sb m Sn IV Cl 1 Br 1 I 1 O n N m C IV As we should expect, an atom with the double capacity, can combine with two of the single capacity, or with one of the double capacity, and so forth. Thus we have compounds of oxygen: Zn n O n Sn IV 2 n Sb a m O 8 11 N 2 1II 3 11 Briefly stated, the quantities of the two elements which com- bine must have equal total combining capacities. Thus Sn IV has the capacity four, and O 2 n has the total capacity of 2 X 2 or 4: Sb 2 m has a total capacity of 2 X 3 (or 6) and so has O 3 n (3x2 = 6). The unit of combining capacity of an atomic weight (or atom) is called a valence. The atomic weights of H and Cl are said to be univalent; those of Zn and O, bivalent; those of Sb and N, trivalent; those of Sn and C, quadrivalent. The highest valence known is eight. Valence and Ionic Charges. Comparison with the formulae of the ions already given will now show that the valence is equal to the number of charges on the corresponding ions: IPCl 1 gives H+ + Cl- and Zn'W gives Zn++ + 2CK Also, of course, the total number of each kind of charges (positive and negative) was equal just as the total valences of each of two constituents of each compound are equal. We shall return, later, to the sup- posed cause of this correspondence between valence and the numbers of charges. Valence of Radicals. What has been said applies to com- pounds of not more than two elements so called, binary compounds. We cannot, by inspection, tell the valences in a compound of three or more elements, like H 2 S0 4 . But, as we VALENCE 183 have seen, all ionogens behave like binary compounds, because they divide into radicals, which move as wholes from one state of com- bination to another. Hence we can assign a valence to the radical S0 4 as a whole. It is evidently bivalent, H 2 I (SO 4 ) 11 , Zn n (S0 4 ) n . Similarly, in K^NOa) 1 , and in ff (NOs) 1 , the N0 3 is clearly uni- valent. Hs^POi) shows P0 4 to be trivalent. Valence also by Displacement. In the foregoing instances, we have learned the valence of an element or radical by studying its combinations. But, clearly, if an element is displaced from combination, atoms of equal total valence must take its place. Thus the action: Zn + 2HC1 -> ZnCl 2 + H 2 shows Zn displacing 2H 1 , and the valence of Zn must therefore be two. We see that this is the case for, on displacing the 2H, it combines with 2C1 1 . Summary. We may now sum up all these facts by saying: The valence of an element is a number representing the capacity of its atomic weight to combine with, or displace, atomic weights of other elements, the unit of such capacity being that of one atomic weight of hydrogen or chlorine. To make a corresponding state- ment for the valence of a radical, we substitute, in the foregoing sentence, the word radical for element, and the word formula- weight for atomic weight. Application in Making Formulae and Equations. We can see at once that the rule of valence will be of great assistance to us in making formulae and equations. Suppose, for example, that we burn a piece of aluminium foil in chlorine, and get the white aluminium chloride. What is its formula? Up to this point, we should simply have looked for it in a book. And if, subsequently, we had required the formulae of the oxide and sul- 184 TEXT BOOK OF CHEMISTRY phate of aluminium, we should have looked these up separately also. But now, all we have to do is to find out the valence of aluminium. Knowing already the valences of Cl 1 and O n and (S0 4 ) n , we have then all the information we require for making the needed formulae. Suppose we know that the atomic weight of aluminium is trivalent Al m (see next section). Making the total valences of each half of the compound alike, we get the formulae: AF'CV, Al 2 ra 3 n , A1 2 III (S0 4 )3 11 . When we know the valences of the elements and radicals, we can make the formula of any required compound. The reader must therefore make a special effort always to learn the valences of each element and radical, and always to use them in making formulae. The reader must also always check every formula he writes from memory, to make sure that it is correct. Thus, if he thinks the formula of zinc nitrate is ZnN0 3 , he must count the valences, Evidently, the correct formula is Zn(NO 3 ) 2 . How to Find out a Valence. To find out the valence of an element, the first time we meet with one of its compounds, take the formula of one simple compound of the element, contain- ing an element of known valence. Thus, what is the valence of thorium? The oxide of this metal forms 99 per cent of the mantel of an ordinary incandescent gas-burner. By referring to any book on chemistry, we find that the formula of this oxide is ThO 2 . The total valence of the oxygen is 2 n or 4. The valence of Th must therefore be four, Th lv . What is the valence of phosphorus in the pentoxide, P 2 5 ? P 2 O 5 n . 2 X x = 5 X 2. Hence x = 5, and the valence of P is five, P v . It is clear, also, that to be able to recall the valence of each element, we must memorize the formula of at least one simple compound of each element. VALENCE 185 Elements with More than One Valence. The rule of valence is not so simple as it has thus far appeared to be. A number of the elements have more than one valence. In other words, the capacity of an atomic weight of such an element may have two (or even more) values, according to the circumstances under which it is combining with other elements. Thus, antimony is usually trivalent, and gives compounds like SbCls, 80263, SbBr 3 . But it can also form compounds in which it is quinquivalent, like SbCls. Similarly, iron forms two complete series of compounds: Bivalent: FeCl 2 , FeO, Fe(OH) 2 , FeSO 4 . Trivalent: FeCl 3 , Fe-A, Fe(OH) 3 , Fe2(SO 4 ) 3 . When an element does give more than one series of compounds, however, we always make a strong point of this fact, so that it may not be overlooked. No simple rule, for telling, in advance, which valence will be used in a given action, can be stated. But the ions Fe++ and Fe 4 "^, for example, have different properties, and are easily recognized in practice. Exceptional Compounds. A few compounds will be met with in which an element shows an exceptional valence. Thus, nitrogen gives two series of compounds of N m and N v . But there are three oxides, N 2 0, NO, and N0 2 , in which the valence of nitrogen seems to be one, two, and four, respectively. However, these are single compounds, not belonging to any series, and are the only compounds of nitrogen showing any of those three valences. Those valences are exceptional. Again, FeO and Fe 2 O 3 belong to the two regular series of com- pounds of iron. But there is the magnetic oxide, Fe 3 O4, where the valence of iron appears not to be a whole number, but 8/3 or 2. In this case the chemist makes the valence regular by sup- posing the magnetic oxide to be a compound of the other two oxides, and writing its formula, FeO,Fe 2 3 . 186 TEXT BOOK OF CHEMISTRY A List of Valences and Charges. The following table contains the valences of some familiar ions and the commonest val- ences of some elements. Many of these elements, however, possess other regular valences, in addition to those shown, so that the list does not pretend to be complete. Univalent. Bivalent. Trivalent. Quadrivalent. Na+ K+ H+ (NH<)+ # Br- F~" (OH)- (N0)- (C10.)- Ca++ Ba-H- Mg+H- Zn++ Pb++ Ni++ Co-H- Mn++ Cu++ (cupric) Fe-H- (ferrous) Hg++ (mercuric) Sn++ (stannous) O (soo- (CO.) C>2 (peroxide) A1+++ Fe+++(ferric) CrH- Sb+++ Bi-H-f (P0 4 ) As (AsH 3 ) B (B 2 3 ) N (NH,, N 2 3 ) P (PH.) Sn-H-H-(stannic) SiO 4 C (CH 4 , C0 2 ) Quinquivalent. N (N 2 5 ) P (P 2 5 ) As (As 2 O 5 ) Sexivalent. S (SO,) Where no charges are indicated, the element, by itself, does not ordinarily form an ion. The Cause of Valence. The number of charges on an ion is equal to its valence. We have recently gained some idea of the reason for this. Electricity is made up of small units of negative electricity, called electrons. The mass of an electron is about one eighteen hundredth of that of a hydrogen atom. A negative charge on an atom of an element means that an electron has been added to the atom. A positive charge means that an electron has been taken from a neutral atom, leaving it positive. Thus an ion of chlorine (Cl~) is equivalent to an atom plus an electron (Cl + ), an ion of hydrogen, an atom minus an electron (H ). When these two ions combine, a neutral molecule of HC1 is formed. Assuming this theory, valence is easily explained. It appears that an atom of chlorine can take up one electron, but not more VALENCE 187 than one, giving a stable structure. H can lose one electron, but not more. Zn, however, can lose two, giving Zn^, and S can gain two, giving S , and Al can lose three A1 + ~ H ' . If Al*" 1 " 1 ", lacking three electrons, combined with only 2C1~, the combination would still have a unit positive charge and would tend to attract another electron (with its attached atom). In this view the atoms of two substances tend to combine in such numbers that the electrons carried by one kind of the atoms (the negative) just suffice to supply the electrons lacking in the other kind (the positive). Chemical combination is a sort of neutralizing operation, and valence is the measure of the excess or deficiency of electrons on each kind of atom or radical. A Suggestion. Having just discussed the conception of valence, we have now considered all the laws of chemical com- position. At this point the reader should pause and review thoroughly the subjects of the first fifteen chapters. The under- standing of the fundamental principles which this retrospect will give will greatly lighten the task of understanding the new and more complex substances we shall have to consider, and the new kinds of reactions and new conceptions we shah 1 encounter, in the chapters immediately following. Exercises. 1. Mark the valences in the Formulae: InCla, V 2 O 5 , OsO 4 , PtCLi. 2. Mark the valences of the radicals in the formulae: Zn(Se04), Al 2 (Te0 4 ) 3 , H 3 (As0 4 ), H(SbO 3 ). 3. If 26 g. of chromium displace 1 g. of hydrogen from hydro- chloric acid, what is the valence of chromium in this displacement (see Table of Atomic Weights) ? 4. Correct the following formulae: ThN0 3 , ThP0 4 , A1(P0 4 ) 2 , LiO, PbF, Bi(N0 3 ) 2 . 5. One gram of a quadrivalent element unites with 0.27 g. of oxygen. What is the atomic weight of the element? CHAPTER XVI CARBON. THERMOCHEMISTRY WE have now covered all the rules regarding the composition of compounds, and we have learned something of the common classes of chemical substances. We next become impatient to know the nature of the kinds of matter composing the bodies of animals and of plants, the materials used in cooking and cleaning, and the coal and illuminating gas we burn. Upon examining wood and coal, fat and starch, soap, sugar, and gasoline, we find that one element is common to all, namely carbon. The same element is found in limestone (calcium carbonate CaC0 3 ) and other rocks. In the present chapter we shall study the forms of the element itself. Its marvelous array of diverse compounds can then be taken up in later chapters. THE FORMS OF CARBON Carbon is known in a variety of forms. Graphite, charcoal, and the diamond are all simply free carbon. We know this be- cause each, when burned in air or oxygen, gives nothing but carbon dioxide gas: C + 2 -> C0 2 . Allotropic Forms. Many substances, like water, may exist in the three states of solid, liquid, and gas. But with most sub- stances three is not the limit. Several solid forms or states, in which the physical properties are as different as are those of diamond and charcoal, are shown by many substances. Such different forms are called allotropes. In the case of carbon the differences are familiar to everyone. The diamond is transparent and crystalline, while charcoal is opaque and non-crystalline. 188 CARL ON 189 The former is, bulk for bulk, three and one-half times as heavy as water, while the latter floats on water and, even when the air is removed from its pores, has a specific gravity less than 2. When burned, 1 g. of carbon gives out differing amounts of heat accord- ing to the form used. Thus diamond gives 7805 calories (p. 193), graphite 7850, and sugar charcoal 8040. The various forms of carbon are all insoluble in ordinary solvents. Molten iron dissolves carbon, however. Diamond. The diamond is distinguished by its natural crystalline form, which often resembles the octohedron (Fig. 39, p. 111). For ornamental purposes it is "cut" by grinding new faces so as to give artificial forms called "brilliants" (Fig. 59) and "rosettes." It is the hardest of familiar substances, and can be scratched or polished only by rubbing with diamond powder. The colorless stones and those with special tints are valuable. The black ("carbonado") and badly colored specimens are less valuable and are used for grinding, for glass-cutting, and on the points of drills. Diamonds are found chiefly in South Africa and Brazil. They are separated from the rock by weathering and washing. They are sold by the carat (1 carat = 4 grains = 205 mgms.*) and the value increases with the size. Thus a first quality, cut stone of 1 carat is worth about $270, one of 2 carats about $340 per carat. The largest known specimen, the Cullinan, weighed 3032 carats before being cut. Graphite is found in nature in Siberia, Cumberland, Brazil, Ceylon and elsewhere. It forms dark grey or black hexagonal tablets, and, when pulverized, it gives slippery scales of microscopic size. Unlike the diamond, it is quite soft, has a specific gravity of 2.3, and conducts electricity. Natural graphite is usually * The new international carat (legal in u. S., 1913) weighs 200 mgms. 190 TEXT BOOK OF CHEMISTRY mixed with foreign matter, and even the purest specimen leaves, when burned, from 2 to 5 per cent of ash. It is called also plum- bago, or black lead. Graphite (Greek, I write), moulded into blocks, is sawn into rods for the cores of "lead" pencils (first used in the 16th century). Clay is added in varying proportions to give different degrees of hardness. Because of its infusibility, it is used to make crucibles. Smeared on a plaster cast (non-conductor), it gives a conducting surface on which metals (copper or silver) can be deposited by electrolysis. A thin layer, used as stove-polish, protects the iron from rusting. In electro-chemical industries it is used for elec- trodes at which chlorine is to be liberated ; all other conductors in- teract chemically with this element and are destroyed. It is employed also as a lubricant, when wooden beams slide upon one another. Large amounts of pure graphite are now manufactured by heating coke with some pitch and a little sand or ferric oxide (Acheson's process). The mixture (3 to 3J tons) is piled (Fig. 60, p. 193) between the electrodes connected with a dynamo, and, on account of its high resistance, becomes strongly heated. The operation is complete in from 24 to 30 hours. Amorphous Forms of Carbon. The non-crystalline or amorphous (Greek, without form) varieties of carbon are numerous. They include wood-charcoal, lampblack, animal charcoal, coal (e.g., bituminous coal and anthracite) and coke. Of these lamp- black (p. 216) contains the greatest amount of free carbon, and bituminous coal practically no carbon, in fact such coal can be almost entirely dissolved by using different solvents. The latter contains the largest proportion of compounds of carbon. Coke is made by heating coal (air excluded) until all the moisture and volatile matter (such as cause burning coal to give smoke) are driven out. Charcoal is made in the same way from wood. These two substances will be discussed after we have learned CARBON 191 something of the chemistry of plants (Chap. XIX), from which they are formed by decomposition. Chemical Properties of Carbon. 1. Carbon unites vigor- ously with oxygen. With an excess of oxygen it forms carbon dioxide, a gas: C + 2 ->C0 2 . With a limited supply of oxygen, it forms carbon monoxide, also a gas: 2C + O 2 -> 2CO. 2. In consequence of this tendency to unite with oxygen, carbon is much used as a reducing agent. Thus, when oxide of copper is heated with pulverized charcoal, carbon dioxide is formed, and the metal is liberated: 2CuO + C -> 2Cu + C0 2 . In the same way the oxides of tin, of lead, and of many other metals may be reduced. Copper, tin, and lead are manufactured from the ores in this way. 3. Carbon unites directly with some elements, particularly with sulphur to form carbon disulphide CS 2 (p. 191) and with certain of the metals. Thus, when dissolved in molten iron, it forms iron carbide Fe 3 C. An interesting case is the union of carbon with hydrogen. The union is ordinarily too slow to be observed. But when the carbon is mixed with pulverized nickel (contact agent) and hydrogen is passed over the mixture at 250, methane CH 4 is formed (99 per cent). At higher temperatures the union is less complete and, at 850, reaches only 1.5 per cent. In the electric arc, however, direct union to form acetylene C 2 H 2 takes place (see van't Hoff's law), because 53,200 calories are required to form it. Dozens of compounds containing only hydrogen and carbon, in different proportions, are known (Chap. XVIII), but all the others are ob- tained by indirect reactions. 192 TEXT BOOK OF CHEMISTRY The valence of carbon is almost always four. This is clearly seen in C^O/ 1 and C^H/. In a few compounds, of which CO is the commonest, carbon is bivalent. Carbon Tetrachloride CC1 4 . This compound is manu- factured by leading dry chlorine into carbon disulphide, in which a little iodine (contact agent) is dissolved: CS 2 + 3C1 2 - CC1 4 + S 2 C1 2 . On distilling the resulting mixture, the carbon tetrachloride CCU (b.-p. 77) passes off and is condensed, while the sulphur mono- chloride S 2 C1 2 (b.-p. 136) remains. The latter compound is purified by separate distillation, and is used in vulcanizing rubber. As now made, carbon tetrachloride costs, wholesale, only 10 cents per pound. Carbon tetrachloride is a colorless liquid. It dissolves fats and tars and other organic compounds, and has the advantage over benzine and gasoline of being non-inflammable. It is there- fore used in taking the grease out of wool, linen cloth, oil-bearing seeds, and bones. "Carbona," sold for removing stains from, and cleaning ("dry cleaning"), clothing, gloves, etc., is benzine (see p. 210) to which sufficient carbon tetrachloride has been added to render the mixture non-inflammable. "Pyrene" fire extin- guishers contain a liquid which is mainly carbon tetrachloride. When the liquid is directed upon burning material, the carbon tetrachloride is vaporized. The vaporization cools the mass by using up the heat, and the vapor at the same time displaces the air and stops the combustion. Carbides and the Electric Furnace. Chemical actions which proceed only at very high temperatures are most economi- cally carried out by using electricity as the source of heat. In such cases the electricity has no electrolytic or other chemical action. There are two types of electric furnaces. In the making THERMOCHEMISTRY 193 of graphite (p. 190) and of carbon disulphide (p. 191), which illustrates one of them, the resistance of the carbon furnishes the occasion for the rise in temperature. Of the same type is the furnace used for making carborundum (SiC, carbide of silicon), of which hundreds of tons are manu- factured annually at Niagara Falls (Acheson's process). The coke and sand (silicon dioxide SiO 2 ) are piled between the ter- minals, and the resistance of the former causes the produc- tion of the heat (Fig. 60) : 3C + Si0 2 -> SiC + 2CO f . _____ Here the carbon reduces the oxide, and combines with the element (Si) as well. The product (SiC) is exceedingly hard, and, after pulverization and mixing with other materials, is moulded into grinding wheels. In the other type of furnace the air between the terminals furnishes the resistance, and the arc (a discharge carried by the badly conducting air and carbon vapor) furnishes the heat. The arc is used in making calcium carbide (CaC 2 ), by heating a mixture of lime (CaO) and coke: CaO + 3C -> CO + CaC 2 . Cold water acts vigorously with calcium carbide, giving acetylene gas C 2 H 2 (see p. 193) and calcium hydroxide (slaked limp.) : - 2H 2 -> Ca(OH) 2 + C 2 H 2 T THERMOCHEMISTRY Heat and Chemical Change. We decompose potassium chlorate because we desire to get oxygen. We burn phosphorus when we require some phosphorus pentoxide. Chemical changes are used for the purpose of making hundreds of needful materials. But chemical changes, in may cases, produce heat, as well as new 194 TEXT BOOK OF CHEMISTRY materials, and it is quite as often the heat, and not the new sub- stance, that we want for commercial purposes. We sometimes use both, but more often one is thrown away. In the steam engine we burn coal, not to obtain carbon dioxide, but heat heat to generate steam and obtain motion and do work of some valuable description. In the gas of gasoline (petrol) engine we burn (explosively) a mixture of air and combustible vapor, and we use the force of the explosion to do work, while the material products of combustion pass into the atmosphere. The heat is therefore to be kept in mind as one of the real and most important products of many chemical changes. Heat and Energy. Heat was formerly thought of as a kind of matter, which simply lacked weight. The heat of the engine is used up in turning the dynamo, and electricity is set in motion. The electrical energy can be turned into light, as in the arc lamp or into heat by passing through a thin wire, as in the tungsten bulb. Light can be turned back into heat by falling on a black body. None of these things have weight, and each can be turned into any of the others. In both of these ways they differ from matter, the forms of which cannot be changed one into another. To mark these differences a separate name is now used, and heat, light, and electricity are called forms of energy. Energy is some- thing that can do work. It resembles matter only in the fact that it can be kept, or changed from one form to another, without loss in its quantity a fact called the principle of the conserva- tion of energy. Hence energy is bought and sold like matter. Internal Energy. The heat in a tungsten lamp arises because electrical energy is being changed. The electricity in a wire comes from the motion and heat of the engine, which are being used up. Since energy cannot be created from nothing, we look for the source of every development of energy in a supply of some other variety. Hence, when heat is given out in a chemical change, THERMOCHEMISTRY 195 we assume that there was energy in some other form in the original matter such as the oxygen and carbon and this form we call chemical or internal energy. It is this energy that gives their value to our coal-beds, our batteries, and our explosives, not the particular materials they may contain. It is the internal energy of coal or coke and oxygen that we use to obtain heat for our houses and factories. We use the internal energy of bread (largely starch) and oxygen, by eating the former and breathing the latter, as a source of the mechanical energy required to enable our bodies to perform greater or smaller amounts of physical labor. The heat or other forms of energy obtainable from chemical actions is therefore of the most vital importance to us as human beings. The Measurement of Heat. The heat produced by a chemical action is measured by carrying out the changes, with weighed amounts of the materials, in a vessel surrounded by water (a calorimeter Greek, heat measurer) . The amount of water is known, and its rise in temperature observed with care. Each gram of water raised one degree represents one calorie of heat developed (see p. 195). Thermochemical Equations. For convenience in chem- istry we calculate, from the data obtained by measurement, the amount of heat (in calories) which would be developed by use of such quantities of the substances as are represented by the formulae in the equation for the change. We can then in- sert the heat as one of the products of the action. Thus, when carbon burns with oxygen to give carbon dioxide, 1 g. of car- bon unites with 2.67 g. of oxygen to give 3.67 g. of carbon dioxide and 8040 calories. Hence, since C stands for an atomic weight (12) of carbon, the amount of heat to be placed in the equation is 12 X 8040 ( = 96,480) calories C + O 2 -* C0 2 + 96,480 cal. 196 TEXT BOOK OF CHEMISTRY Many actions absorb heat (are endothermal), instead of liberat- ing it, as does the burning of carbon (exothermal). Thus, the dissociation of 36 g. of water vapor to free hydrogen and oxygen absorbs 4 X 28,800 = 115,200 cal. This absorption is indicated by the negative sign preceding the number of calories: 2H 2 O -> 2H 2 + 2 - 115,200 cal. When the action is reversible, the heat absorbed when it goes in one direction is liberated when the action proceeds in the other direction. 2H 2 + 2 5 2H 2 O + 115,200 cal. An action which absorbs heat can take place only if heat or some other form of energy is furnished. Thus, the decompo- sition of hydrogen chloride, by electrolysis of hydrochloric acid (p. 55), consumes electrical energy equivalent in amount to the heat given out when hydrogen and chlorine unite to form hy- drogen chloride in solution: H 2 + C1 2 + Aq - 2HC1, Aq + 78,800 cal. Exercises. 1. (a) What physical property of graphite enables it to cover the surface of a stove so effectively? (b) How does " polishing" with a brush contribute to the result? (c) Why not use paint on a stove ? (d) Explain why graphite can be used as a lubricant. 2. If a metal formed the positive electrode (anode) in electrolys- ing sodium chloride solution, what chemical change might it undergo (p. 54), and which metals would be least rapidly attacked? What objection is there to using the latter metals in practice? 3. When one cubic meter of oxygen acts upon carbon, what volumes (at the same temperature and pressure): (a) of carbon dioxide; (b) of carbon monoxide can be obtained? 4. Make the equation: (a) for the formation of methane by union of carbon and hydrogen; (b) for the reduction of stannic oxide (Sn0 2 ) by carbon. THERMOCHEMISTRY 197 5. What form of energy is delivered, and paid for as such, in all cities ? 6. Make thermochemical equations for the combustion of diamond and graphite. 7. From the amounts of heat given out by the different forms of carbon, when burned, what do you infer as to the relative amount of energy in each form? Which form contains the most energy? CHAPTER XVII THE OXIDES OF CARBON CARBON unites with oxygen in two different proportions, forming carbon dioxide C02, and carbon monoxide CO. These oxides differ very markedly in properties. They both have im- portant, although very different, uses. Carbon dioxide is the gas which escapes from effervescent beverages (soda water, beer, and champagne) and immense quantities of it are used in making the carbonates of sodium. Carbon monoxide is less familiar, but very important as a constituent of fuel and illuminating gases, and as the substance directly responsible for the reduction from the ores of all the iron and steel so freely used for rails, structural iron, and innumerable other articles. CARBON DIOXIDE Occurrence. Carbon dioxide is found in nature issuing from the ground, especially in volcanic neighborhoods, and dissolved in effervescing natural waters, such as Saratoga and Vichy. It is found in the air (4 liters in every 10,000 liters of air) and in the breath (37 liters per 1000 liters). Preparation. 1. Carbon dioxide is most easily prepared in the laboratory by the action of an acid such as hydrochloric acid upon a natural carbonate like calcium carbonate (marble or lime- stone). The action occurs in two stages. The first is a double decomposition, such as all acids and salts exhibit when brought in contact with one another : CaC CaCV+ H 2 C0 3 . 198 CARBON DIOXIDE 199 The calcium chloride CaQ 2 remains dissolved in the water con- tained in the hydrochloric acid. The carbonic acid H 2 CO 3 is unstable, however, and immediately dissociates into water, which remains and carbon dioxide gas which escapes: H 2 CO 3 ^ H 2 O + CO 2 T The apparatus used is similar to that employed in making chlorine (Fig. 47, p. 135). 2. For commercial purposes the carbon dioxide is either used as it is produced, or else it is compressed into wrought-iron cylin- ders and shipped in the form of a liquid. Three sources of such commercial carbon dioxide are in use: When carbon, for example, in the form of coke, is burned with a plentiful supply of air, all the carbon is converted into carbon dioxide. C + 2 - C0 2 . Since, however, there are four volumes of nitrogen to one of oxygen in the air, this carbon dioxide is diluted with nitrogen in the same proportion. The gases must therefore be separated by leading them through potassium carbonate solution, which absorbs the carbon dioxide: C0 2 + H 2 <=> H 2 C0 3 + K 2 C0 3 ^ 2KHC0 3 . The bicarbonate of potassium KHC0 3 thus produced is subse- quently decomposed by heating. 3. Calcium carbonate (limestone) CaC0 3 , or, more easily, mag- nesium carbonate (magnesite) MgCO 3 , may be decomposed by heating in a kiln: CaCO 3 -> CaO + CO 2 T MgC0 3 -> MgO + C0 2 T In the former case the quicklime (CaO), which is formed at the same time, is a valuable product also. 4. Carbon dioxide is formed in fermentation, and so is collected from the vats in which beer is brewed (see p. 223). 200 TEXT BOOK OF CHEMISTRY Physical Properties. Carbon dioxide is a colorless, odorless, almost tasteless gas. As the molecular weight (CO2 = 44) shows, it is one-half heavier than air. Its greater specific gravity may easily be shown by pouring it from one jar into another, or into a beaker placed on one pan of a balance with an equipoise of shot on the other pan. It is much more soluble in water than is air. One volume of water at 15 dissolves an equal volume of the gas. At two atmospheres pressure, two volumes are dissolved, at three atmospheres, three volumes. Pure water, charged at 3 or 4 atmospheres pressure is known as soda water. Effervescent waters, such as Selters and Vichy, contain dissolved salts in addi- tion. The gas can be liquefied at any temperature below 31.35. At 20 the pressure required is 60 atmospheres and this is therefore the pressure in a cylinder of liquid carbon dioxide at that temperature. To withstand the pressure very massive cylinders are required, and they weigh, when empty, about twice as much as does the liquid they will hold when full. When liquid carbon dioxide is allowed to run from a cylinder into a cloth bag (non-conductor of heat), the rapid evapora- tion of a part of the liquid consumes so much heat that the rest of the liquid freezes to a snow-like mass of solid carbon dioxide. In the laboratory this solid is used as a cooling agent, being mixed with ether or benzine to secure closer contact with the object to be cooled. Chemical Properties. Stability. 1. Carbon dioxide is very stable (only 7.5 per cent dissociated at 2000), and so, although it contains much oxygen, it will not support combustion. Mag- nesium or aluminium powder (see activity list, Appendix V), how- ever, will burn when placed on a cake of solid carbon dioxide and set on fire with burning magnesium ribbon; 2Mg + C0 2 -> 2MgO + C. CARBON DIOXIDE 201 The gas extinguishes burning wood, oil, or candles, and 15 to 16 per cent of it in air is sufficient to extinguish ordinary com- bustibles. For this reason some fire extinguishers contain a dilute solution of bicarbonate of sodium (NaHCO 3 , p. 353) and sulphuric acid. 2NaHCO 3 + H 2 SO 4 ^Na^SC, + 2H 2 CO 3 - 2H 2 + 2CO 2 T 2Na+ + 2HCO 8 - + 2H+ + SO 4 = -f 2Na+ + SO 4 = + 2H 2 C0 3 ->2H 2 O + 2C0 2 T When the instrument is inverted, these materials are mixed, and water and carbon dioxide are forced out by the pressure of the gas. Disused wells and excavations should be tested with a lighted candle before being entered, since a proportion of carbon dioxide, slightly larger than the 15 per cent which extinguishes a candle, will cause suffocation. Chemical Properties Carbonic Acid. 2. Carbon di- oxide, when dissolved in water, combines in part to form carbonic acid: The gas is therefore often called the anhydride (Greek, without water) of carbonic acid. The solution has all the properties of an acid, although, as the acid is very little ionized (p. 170), it exhibits them rather feebly. It tastes slightly sour, turns blue litmus faintly red, and neutralizes bases. The last action is easily shown by shaking the gas with limewater (solution of calcium hydroxide, a base) : Ca(OH) 2 + H 2 C0 3 -> CaC0 3 J + 2H 2 O. The carbonate of calcium is precipitated and the liquid becomes milky in appearance. This action is used by sugar refiners for removing the lime employed in purifying the sugar. Manu- facturers of white lead (carbonate of lead) also employ carbon dioxide, because of its entering into double decomposition to give 202 TEXT BOOK OF CHEMISTRY carbonates. The same property is utilized in making bicarbonate of sodium and washing soda (carbonate of sodium, p. 353). Since the molecule of carbonic acid (H^COs) contains two atoms of hydrogen, either one or both of these atoms may be replaced by a metal the acid is dibasic. Thus we may have sodium car- bonate Na 2 C03, or its hydrate, washing soda Na 2 C0 3 , 10H 2 0, and also sodium bicarbonate (baking soda, sodium-hydrogen carbonate) NaHCO 3 (see pp. 352-354). 3. The most marvellous chemical action into which carbon dioxide enters is perhaps the most useful to mankind, and at the same time the one least understood. This is the action by which plants use the gas as food. This action is so important that the last section of the present chapter is devoted to it. CARBON MONOXIDE Preparation. Carbon monoxide CO is most easily prepared in the laboratory by heating formic acid (or sodium formate, a white crystalline solid) with concentrated sulphuric acid. The latter combines with the water, but is not otherwise changed : HC0 2 H -> CO + H 2 0. When coke, or any form of carbon burns with a limited supply of air, or oxygen, the same gas is produced: 2C + 2 -+ 2CO. The gas therefore rises from the surface of a coal fire, sometimes escaping unburned, but often burning with a blue flame above the coal. Producer Gas and Water Gas. When air is used with the coke, the mixture of carbon monoxide (39 per cent) with nitrogen (60 per cent) obtained is called producer gas. It is combustible, and is used in industrial establishments for heating and to drive gas engines for power. Commercially, large amounts of carbon monoxide, mixed with CAKBON MONOXIDE 203 hydrogen (water gas), are manufactured by blowing steam over white hot coke or anthracite : C + H 2 O -> CO + H 2 - 28,300 cal. The coke is first set on fire in a brick-lined cylindrical structure and brought to vigorous combustion by blowing in air for ten minutes. Then steam is substituted for the air. The interaction, as the equation shows, takes place with ab- sorption of heat. Hence, at the end of about five minutes, the coke becomes cooler. It is then necessary to turn the steam off and to turn the air on again, and so on alternately. The mixture of carbon monoxide (40 to 50 per cent) and hydrogen (45 to 50 per cent), containing also some carbon dioxide (4 to 7 per cent), nitrogen (4 to 5 per cent), and oxygen (1 per cent), is known as water gas. It is almost wholly combustible, burning with a blue flame, and is used as a source of heat and, by driving internal com- bustion engines, to furnish power. It is used also in manufactur- ing illuminating gas (see p. 213). If both air and steam are driven together over the burning coke, the air enables the coke to burn continuously, and a fuel gas which is a cross between producer gas and water gas is obtained. Fuel gases are employed on a large scale in steel works, and other industrial plants. They give a uniform and easily regulated heat, they leave no ash, and their use involves no labor for stoking. Physical Properties. Carbon monoxide is a colorless, odorless, and tasteless gas. It is a little lighter than air (mol. wt. 28), and is very slightly soluble in water. It is one of the gases which are most difficult to liquefy and its boiling-point, when liquid, is 190, close to that of liquid air. Chemical Properties. When set on fire, the gas burns in air or oxygen with a blue flame. Carbon dioxide is formed, and the presence of the latter may be shown with lime-water (p. 201) ; 2CO + 2 -> 2C0 2 . 204 TEXT BOOK OF CHEMISTRY On account of this property, carbon monoxide reduces the oxides of the less active metals, such as those of iron and of the metals below iron in the order of activity (Appendix V) . Commercially, the ores of iron are reduced by this gas (essentially producer gas) in the blast furnace. The oxides of the metals above iron are not reduced. Physiological Properties. The gas is an active poison, and 1 volume in 100,000 volumes of air produces symptoms of poison- ing, while one volume in 750 to 800 volumes produces death in about thirty minutes. The gas combines with the haemoglobin of the blood corpuscles, forming a stable compound, and thus preventing the absorption of oxygen by the blood (p. 34). This gas is the chief poisonous substance in illuminating gas. The poisonous effect of tobacco smoke, particularly when inhaled, is due to the carbon monoxide produced by the necessarily incom- plete combustion. When the "damper of a stove is almost closed, the carbon mon- oxide rising from the surface of the fire (p. 191) does not escape freely into the chimney, and is apt to find its way, through the crevices, out into the room. In such circumstances, if the room is not well ventilated, a poisonous proportion of the gas may easily be reached. CARBON DIOXIDE AS PLANT FOOD How the Plant Feeds. The walls of the cells which form the frame-work of a plant are made of cellulose (CeHioOs)^ In the cells, especially in certain parts of the plant, granules of starch (CeHioOs)* are found. These substances differ in properties, al- though they have the same composition. The plant contains also protoplasm and proteins. Now all these substances contain carbon, hydrogen, and oxygen, and plant food must furnish these elements. Hence, in addition to large quantities of water ascend- ing from the soil through the roots and stem, and sufficient amounts of compounds of nitrogen, potassium, and phosphorus, all plants CARBON DIOXIDE AS PLANT FOOD 205 require an abundant supply of carbon in absorbable form. This carbon is practically all taken up by plants in the form of atmos- pheric carbon dioxide. It is admitted through minute openings (stomates), situated mainly in the surface of the leaves. The Reaction Involved. Comparison of the formulae of carbon dioxide CO 2 and of any plant substance, like starch (CeHioOs)*, shows at once that the latter contains a far smaller proportion of oxygen, relatively to the amount of carbon, than does the former. Hence, during the digestion or assimilation of the carbon dioxide by the plant, this compound must be reduced. In point of fact, the chlorophyll (green matter) and proto- plasm in the leaves act upon the carbon dioxide, causing oxygen to be liberated : 6C0 2 + 5H 2 -> CeHioOs + 6O 2 T FIG. 61. This action goes on only in the sun- light. The steps by which sugar, starch, and cellulose are manufactured by the plant out of water and carbon dioxide, are not well understood. But the liberation of the oxygen is easily shown by placing a green plant under water in a jar, and setting the jar in the sunlight (Fig. 61). Bubbles of gas appear on the leaves, grow larger, and then detach themselves and rise to the top. The gas relights a glowing splinter of wood, and is pure oxygen. The Thermochemistry of the Reaction. In the combina- tion of carbon and oxygen, during combustion of wood or coal, much heat is liberated. Hence, when oxygen is taken out of carbon dioxide again, heat or energy in some form must be supplied. When this takes place in a plant, the energy is evidently furnished 206 TEXT BOOK OF CHEMISTRY by the sunlight, for the action proceeds more slowly in the shade, and ceases in the dark. The energy required can be measured, and may be expressed in calories. The energy required to produce one formula-weight of cellulose (C 6 Hi O 5 = 6 X 12 + 10 X 1 + 5 X 16 = 162 g.) is 671,000 calories. The whole may be represented in a rough equation, in which the unknown intermediate steps are left out, and only the starting substances and the final products are shown : 6CO 2 + 5H 2 O + 671,000 cal. --* C 6 Hi 5 + 6O 2 . Source of the World's Energy. The energy that does the world's work comes mainly from two sources, namely, water power and the combustion of wood, or of coal (which is fossil wood). The water comes from vapor, generated by the sun's heat, con- densed as rain, and collected in lakes or reservoirs. The source of the energy of coal or wood is a little less obvious. When wood (which is largely cellulose) burns, it gives carbon dioxide, water, and heat. In fact, its combustion is represented in the above equation, when the equation is read backwards. Thus the sunlight, working through the machinery of the plant, takes the carbon dioxide and water, furnishes the energy (as light), and gives us wood and oxygen. And the wood and oxygen, when burned, give us back the original substances, and the equivalent of the original energy in the form of heat. Hence, our other main source of energy turns out to be the same as the first the sun's rays although the route by which the energy comes to us is a little less direct. If, instead of burning the starch of the plant, we consume it as food, it goes through a series of changes instead of only one. But the end products are the same, namely, carbon dioxide and mois- ture issuing from our lungs, and heat and other forms of energy such as are developed in living organisms. Thus, whether we use our muscles, a steam engine, or a waterwheel to do work, sunlight is in each case the ultimate source of energy employed. CAKBON DIOXIDE AS PLANT FOOD 207 Exercises. 1. Make equations for: (a) the action of sulphuric acid upon calcium carbonate; (b) carbon dioxide on sodium hydroxide solution (p. 201) ; (c) the burning of aluminium in car- bon dioxide. 2. From the fact that the molecular weight of carbon dioxide is 44, how do we infer that it is one-half heavier than air? 3. Why does soda water remain quiescent in the closed bottle, and why does it effervesce when the bottle is opened? 4. What may be the sources of carbon dioxide in disused cellars and excavations? 5. Name the variety of chemical change (p. 127) to which belongs the reaction shown in each equation in this chapter. 6. Assuming that air contains oxygen and nitrogen in the pro- portion of 1 : 4 by volume, what are the theoretical proportions of carbon monoxide and nitrogen in producer gas? 7. (a) What volume of water gas is produced from each liter of steam, and (b) what is the proportion of the component gases in the product? (c) What impurities should you expect to find in water gas? (d) How should you attempt to separate the com- ponents of water gas ? 8. Why is water gas an especially valuable source of heat when high temperatures are required. 9. Make a list of metals the oxides of which would be reduced by carbon monoxide. 10. In what relative volumes are carbon dioxide used and oxygen produced by a plant? CHAPTER XVIII COMPOUNDS RELATED TO PETROLEUM. FLAME PETROLEUM introduces us to the compounds of carbon and hydrogen, called by chemists the hydrocarbons. Many scores of these are known, and a number occur in nature. Thus, the natural oil, petroleum, is a mixture of hydrocarbons. To under- stand its treatment and uses, we must consider first the nature of its components. The Hydrocarbons. The following list contains the names, formulae, and boiling-points of seven of the simplest hydrocarbons and of two others. Methane CH 4 b.-p. -164 Hexane C 6 H 14 b.-p. 71 Ethane C 2 H 6 -89.5 Heptane C 7 Hi 6 99 Propane C 3 H 8 -37 Hexadecane Ci 6 H 34 287.5 Butane C 4 Hi + 1 Hexadecane Ci 6 H 34 m.-p. 18 Pentane C 6 Hi 2 35 Pentatriacontane C 3& H 7 2 m.-p. 74.7 After the first four, the names are based on the Greek numerals corresponding to the number of carbon atoms in the molecule. Thus heptane is followed by octane CgHi 8 , nonane C 9 H 20 , decane CioH 2 2, and so forth. On comparing the formulae, we observe that rJ in each, the number of units of hydrogen is equal to twice the T number of carbon units plus two. The general formula is therefore |t C n H 2 nf.2. The series affords a striking illustration of the law of multiple proportions. We note, further, that the first four are gases at the ordinary temperature. The members of the series from pentane to pentadecane (Ci 5 H 32 ) are liquid under ordinary conditions. From hexadecane onwards they are solids, with higher and higher melting-points. Even this long series, the numbers of which have been studied 208 COMPOUNDS RELATED TO PETROLEUM 209 up at least as far as hexacontane CeoHua (m.-p. 102), does not include all the known hydrocarbons. Thus, a compound, to the presence of which illuminating gas largely owes its luminous flame, is ethylene C 2 H 4 . Starting with this compound, another series (C n H 2n ) is known, of which the molecules contain 2 atoms of hydrogen less than the corresponding compounds of the first series. There are still other series, such as that beginning with acetylene C2H 2 (C n H 2 n-2), in which the proportion of hydrogen is relatively still smaller. The ethylene, acetylene, and other series, in which the propor- tion of hydrogen is less than the maximum, are spoken of as Tin- saturated hydrocarbons because the full valence of the carbon is not in use, and these compounds unite with hydrogen, chlorine, and other elements rather readily. On the other hand, the mem- bers of the first series are called the saturated hydrocarbons. These can take up other elements only by giving up a part of their hydrogen to make room for the new element. All these hydrocarbons are mutually soluble, so that mixtures can be made in any proportion, and using any number of members of the different series. Petroleum. Petroleum is a thick, greenish-brown oil. When borings are made into the oil-bearing strata, the oil either gushes up, or is pumped to the surface. Wells of this kind are in opera- tion at Baku on the Caspian, in India, in Japan, in Ontario, and in Ohio, Pennsylvania, California, and Oklahoma in North Am- erica. In the United States hundreds of miles of pipe-lines are used to transport the oil, with the aid of force pumps, to the re- fineries, and in 1912 nearly 220 million barrels (42 gal. each) were produced. The world's production in 1912 was 350 million barrels. Oil Refining. The natural oil is a complex mixture, and the components are partially separated in order to get products suit- able for various purposes. The separation is carried out by dis- 210 TEXT BOOK OF CHEMISTRY tillation. The components of lower boiling-point come off first. Then the remaining oil is chilled and a quantity of the solid mem- bers of the series (C22H 46 to C28H 58 ) crystallizes in flakes (solid paraffin) and is separated by filtration. The unsolidified oil is distilled further to get the components of intermediate boiling- point. The residue is used for lubricants and for fuel. The various products are still mixtures, but contain only compounds lying close together in the series. Some of the products, together with their main components and uses, are as follows: Name. Main Components. B.-P. Uses. Petroleum ether... Gasoline . Pentane, hexane Hexane heptane 40- 70 70- 90 Solvent, gas-making Solvent fuel Naphtha Heptane, octane 80-120 Solvent, fuel Benzine . . . Octane, nonane 120-150 Solvent Kerosene Decane-hexadecane 150-300 Illuminating oil Vaseline, C 2 2H 46 to C 2 3H 48 , is separated in some refineries. Solid paraffin is employed for waterproofing paper, as an ingredient in candles, and in making chewing gum. Asphalt, a natural mixture of the solid hydrocarbons, found particularly in Trinidad, is used in road-making. Natural Gas: Methane CH 4 . Natural gas is obtained from wells, tapping strata close to those which contain petroleum, and in the same localities. It often issues under very high pressure. It owes its combustibility to its chief component (over 90 per cent), methane CH 4 . It is largely used as a fuel in the regions in which it is found and, in the United States, the annual value of the gas so consumed is estimated at nearly $73,000,000 (1910). The same gas issues from many coal seams (" fire-damp")? and forms ex- plosive mixtures with the air of mines. It rises to the surface when stagnant pools containing decomposing vegetable matter are stirred (" marsh-gas"). COMPOUNDS RELATED TO PETROLEUM 211 Acetylene C2H 2 . A mixture, containing acetylene, is formed when any hydrocarbon is heated strongly (p. 212), air being excluded. Pure acetylene is prepared by the action of water on calcium carbide (p. 193) : CaC 2 + 2H 2 -* Ca(OH) 2 + C 2 H 2 T Calcium hydroxide (slaked lime) remains. The gas burns with a flame even more luminous than that of ethylene. It is there- fore made in generators by the foregoing action for use on auto- mobiles and for lighting buildings remote from a public supply of illuminating gas. Acetylene tanks, which are also in use, contain acetylene dissolved, under high pressure, in acetone. The Acetylene Blowpipe or Torch. Acetylene decomposes, when heated, with liberation of heat: C 2 H 2 -* 2C + H 2 + 58,100 cal. When acetylene burns with oxygen, therefore, 2C 2 H 2 + 5O 2 -+ 4CO 2 + 2H 2 O, we obtain not only the heat due to the combustion (p. 189) of the carbon (4 X 96,480 cal.) and of the hydrogen (4 X 28,800 cal.) but also the heat due to the decomposition of the gas (2 X 58,100 cal.). The temperature of the flame is, therefore, the highest that can be reached by the combustion of any easily obtainable gases. The oxy-acetylene flame, produced by means of a suitable burner (Fig. 29, p. 56), the gases being furnished from small, portable tanks, is now used for cutting metals. Such a flame will melt its way through a 6-inch shaft of steel, or a heavy steel plate several feet wide, in less than one minute, cutting the object in two. Steel buildings have been taken apart rapidly by this device. Blau gas and oil gas, mixtures of hydrocarbons made by " crack- ing" (see next two sections) heavy oils, are now largely displacing acetylene for uses like those just mentioned. They give flames 212 TEXT BOOK OF CHEMISTRY which are almost as effective, and are more easily controlled. Even the oxy-hydrogen torch is remarkably efficient, when applied to the same purposes. Chemical Properties of the Hydrocarbons. The hydro- carbons, whether pure or in solution, show no conductivity for electricity. They have none of the chemical properties of acids, bases, or salts, and therefore do not enter into double decom- positions with substances of these classes. The saturated hy- drocarbons are in fact quite indifferent to the presence of most chemical reagents. The unsaturated hydrocarbons act chiefly by combining with hydrogen, the halogens, and some acids: C 2 H 4 + C1 2 -^ C 2 H 4 C1 2 (ethylene chloride). All the hydrocarbons burn with oxygen or air to form carbon dioxide and water: CH 4 + 20 2 4 C0 2 + 2H 2 0. C 7 H 16 + H0 2 - 7CO 2 + 8H 2 O. The water can be shown by its condensation on a cold vessel held over the flame. The carbon dioxide gives a precipitate of calcium carbonate (p. 201) when the gases rising from the flame are drawn through lime-water. All the hydrocarbons, when heated strongly (air excluded), decompose or crack. They usually lose a part of their hydrogen and become unsaturated. These of high molecular weight break up to give a mixture of hydrocarbons of low molecular weight. Ethylene C 2 H 4 , for example, is produced in large amounts by heating the higher members of the series to a red heat. On the other hand, the lower members of the series, when heated, often give compounds of higher molecular weight. Thus, methane gives ethylene and acetylene, along with hydrogen, 2CH 4 - C 2 H 4 + 2H 2 . 2CH 4 > C 2 H 2 4~ 3H 2 . COMPOUNDS RELATED TO PETROLEUM 213 At a white heat all the hydrocarbons decompose into hydrogen and free carbon: CH4 -> C + 2H 2 . Carburetted Water Gas. To fit water gas, essentially H 2 + CO (p. 203), which burns with a pale blue flame, for public service as an illuminating gas, unsaturated hydrocarbons, and particularly ethylene CzHA, which burn with a highly luminous flame, must be added. The water gas is passed through a tower, filled with strongly heated brick on which oil is continually sprayed. Mixed with the vapor of the oil, the gas goes into the "super- heater" where, at a higher temperature, the decomposition into unsaturated hydrocarbons (cracking) takes place. The gas is then cooled and washed to remove the condensible hydrocarbons, which would otherwise collect in the service pipes, with resulting waste of combustible material as well as obstruction in the de- livery of the gas. A typical carburetted water gas has the com- position: Illuminants (largely ethylene) 16.6 per cent; heating gases methane 19.8 per cent, hydrogen 32.1 per cent, carbon monoxide 26.1 per cent, impurities (nitrogen and carbon dioxide) 5=4 per cent. Sources of Gasoline. The demand for gasoline (petrol) for use in automobiles and motor launches has recently become greater than the natural supply of the more volatile components of the natural oil. Two new sources of gasoline have been found. One plan is to crack the heavier hydrocarbons by heat in such a manner as to get a fair yield of the saturated compounds boiling between 70 and 90. The other is to apply pumps to the disused gas wells, from which natural gas no longer issues; and so, by creating a vacuum, to promote the evaporation of the volatile part of the liquids remaining in the strata. The vapor so obtained in then condensed. 214 TEXT BOOK OF CHEMISTRY FLAME We have encountered a variety of flames, from the simple one of hydrogen burning in air to the more complicated case of the luminous flame of ethylene or acetylene. The subject will now repay a somewhat closer study. The Simple Flame. The flame of hydrogen (giving water), or of carbon monoxide (forming carbon dioxide) is very simple in structure (Fig. 62). We find that there is a tapering column of unburnt gas in the interior, surrounded by a layer of hot /\ gas the flame itself. The flame is therefore a hollow cone. That the flame is hollow is easily shown by hold- ing a wooden match across it. The match is charred at the two points at which it crosses the flame, and remains unheated in the middle. These flames are simple, be- cause only one chemical change occurs in them. The .rIG. u2. flames are rather large, because sufficient oxygen to burn all the gas does not reach the latter at once, and the gas travels upwards and diffuses outwards a certain distance before being all consumed. If oxygen is substituted for air, by lowering the jet into a jar of that gas, the flame becomes much smaller. In the absence of atmospheric nitrogen, there is now five times as much oxygen within a given range of the center of the jet as before. This chemical union, like any other, proceeds more rapidly with an increase in the concentration of the interacting substances (p. 35). It is therefore completed before the gas has time to diffuse very far from the opening of the jet. The Candle Flame. A candle is made of a mixture of paraffin and stearic acid (a compound of carbon, hydrogen with some oxygen, made from fat). When it burns, the whole phenomenon is vastly more complicated than the burning of hydrogen. The FLAME 215 following are some of the stages in the process, which is operated by the flame's own heat. To start with, the wax is melted and ascends the wick by capillary action. This is merely a physical phenomenon. Then the chemical changes begin. (1) The melted compounds of carbon are decomposed by the heat (cracked, p. 212), being turned into more volatile compounds and gases which occupy the central hollow of the flame. (2) The compounds forming the gases and vapors are further decomposed at a white heat, giving free carbon and hydrogen (p. 213). (3) All the materials finally reach a sufficient supply of oxygen and are burned to water and carbon dioxide. There are thus three chemical changes, each of which takes place in a definite region that can be observed by the eye (Fig. 63). The formation of the gases from the melted wax (without gas, there would be no flame) takes place in the dark central region where there is no oxygen. The carbon is set free and glows brilliantly in the luminous cone that surrounds the gas and extends far above it. The final combustion occurs in a fainter cone of flame covering the whole exterior. That there is unburnt gas (produced by decomposition of the wax) in the center is easily shown by inserting a narrow tube, through which some of the gas will ascend. The free carbon in the luminous zone will show its presence by blackening a cold dish placed across the flame. Ordinary Illuminating Gas. The candle is a miniature gas factory. It makes its own gas, and burns it as well. It is more convenient and economical, however, to leave the manu- facture of the gas to a public corporation, which distributes it in pipes. In large cities the gas is usually carburetted water gas (p. 213). In many cases, however, coal gas, made by distilling coal, is the kind delivered. In the United States, in 1909, the quantities delivered by public service companies, expressed in 216 TEXT BOOK OF CHEMISTRY billions of cubic feet were: water-gas, 80; coal-gas, 20; mixtures of both, 40. The luminosity of the flame depends upon the complete de- composition of the hydrocarbons (particularly, the ethylene) in the gas. This decomposition with liberation of carbon, is brought about, as in the candle, by the heat of the surrounding mantel of flame. Lampblack. When an iron vessel, cooled by a stream of water circulating through it, is suspended in the flame of natural gas or burning petroleum, the carbon (soot) is deposited on the vessel. By rotating the latter, the soot can be continuously scraped off by a stationary piece of metal. The product, lampblack, being very finely divided carbon, is used in making printers' ink, India ink, and black varnish. Non-luminous Gas Flames. When the gas is to be used for heating, the complete combustion of the gas, without any intermediate liberation of free carbon, is desirable. This is achieved in the Bunsen burner (Fig. 64) by admitting air at the bottom of the burner, in such a way that the air FIG. 64. mixes with the gas before the latter reaches the flame. The air cools the middle zone of the flame, so that at this point the temperature required for dissociating the ethylene, and liberating carbon, is not reached. The oxygen in the air plays no part mixing carbon dioxide or pure nitrogen with the gas has exactly the same effect. A flame of this kind is non-luminous. Although the middle zone of the non-luminous flame is cooler than that of the luminous flame, the average temperature of the flame as a whole is higher. This is the case because the same total amount of heat is liberated in both cases, but the non- luminous flame as a whole is smaller in size. FLAME 217 The Bunsen type of burner, placed in a horizontal position (Fig. 65) is used in the ordinary gas cooking range. As with the Bunsen burner, some care is required to get good results. The holes which admit the air to the mixer must be kept clear of obstructions, as otherwise luminous flames are produced, smoke and soot are formed, and less heat is generated. The size of the openings must be adjusted, as the ad- mission of too much air causes the flame to flash down the burner, and set fire to the gas within the mixer. V Flames with Incandescent Mantels. When gas is burned in a Bunsen burner, a bright light may still be obtained from the flame. This is managed by suspending in the flame a structure (" mantel") made of the oxides of thorium (99 per cent) and of cerium (1 per cent). That these oxides could cohere well enough to withstand ordi- nary treatment, could give out a brilliant, white light when thus heated, and could be manufactured at a price consistent with use in any household, was discovered by Auer von Welsbach. For the same amount of gas passing through the meter the Welsbach lamp gives four times as much light as an ordinary burner. It does this because almost no heat-rays are emitted. Exercises. 1. When vegetable matter decays in the air the carbon it contains is finally all turned into carbon dioxide. When the same matter decays under water, it gives methane (p. 210). Explain the difference in the result. 2. (a) Given a flame of hydrogen burning in a jar of air, what would be the effect on the flame of lowering the pressure of the air by means of an air pump? (b) What would be the effect on FIG. 65. 218 TEXT BOOK OF CHEMISTRY the average temperature of the flame? (c) How about the heat produced by burning 1 g. of hydrogen in each case? (d) What differences would be observed in using an alcohol lamp at the bottom and on the top of a high mountain? 3. (a) In the candle or gas flame, what is the source of the light? (b) Why does such a flame become smoky when placed in a draft? 4. What is the density (air = 1) of (a) methane, (b) ethylene? CHAPTER XIX STARCH AND COMPOUNDS THAT CAN BE MADE THEREFROM WE have already seen (p. 206) how the plant takes carbon dioxide from the air, and water from the ground, and manufac- tures starch C 6 Hi 05 out of them. All the higher plants carry on this important process on a huge scale. Of these plants many, such as wheat, oats, maize (corn) and potatoes, are cultivated for use as food, and the starch is one of the chief foodstuffs that they contain. Furthermore, from starch we can readily make sugar, alcohol, vinegar and other familiar substances. Starch, in fact, introduces us to a much greater variety of organic compounds than does petroleum. Starch (CeHioOs)*. Starch is found in plants in little colorless granules of various rounded shapes (Fig. 66) which may readily be seen under the microscope. These granules are massed in large numbers in the ears of wheat and oats, in the tubers of potatoes, in the grains of maize (corn) and in peas and beans. Even in the leaves they can be seen, ^^ immediately after the plant has been exposed to sunlight. They gradually disappear from the leaves in the dark. They can be recognized, not only by their appear- ance, but, without a microscope, by the iodine test. When a drop of a potassium iodide solution, rendered brown by the addi- tion of a little free iodide, is placed on the leaf or other part of the plant, the granules of starch become blue while the other parts are not affected. 219 220 TEXT BOOK OF CHEMISTRY Preparation of Starch. If flour, which is made by grinding wheat, and is three-fourths starch, is placed in a muslin bag and kneaded under water, the granules of starch are washed out and render the water milky. After a time the granules settle and the water can be poured off. Starch is manufactured by washing disintegrated potatoes (in Europe) or maize (in America) on sieves, and collecting and drying the white powder deposited in the water used for the washing. Starch is not soluble in water. If it be boiled with water, how- ever, the granules swell and break, and the starch becomes finely diffused through the water, forming a clear liquid. With little water, a sort of transparent jelly is produced. When the liquid is poured through a filter, a large part of the starch goes through the paper as if it were truly dissolved. Such a liquid is called a colloidal suspension. Imitation solutions like this are constantly met with in using complex organic compounds such as enter into jellies, glues, soaps, and the juices of the bodies of animals. Even inorganic substances, of the insoluble class, give such suspensions. A description of their peculiarities must be noticed under soap (p. 378). The colloidal suspension of starch turns blue throughout when the iodine solution is added to it. The colloidal suspension of starch is used in the laundry, for stiffening white goods. Glucose is manufactured from it. Glucose C 6 Hi2O 6 /rom Starch. When starch is boiled with water, to which a few drops of an acid (contact agent) such as hydrochloric acid have been added, the liquid, after neutraliza- tion of the acid, is found to be sweet in taste. A kind of sugar, glucose CeH^Oe, can be obtained in crystals by evaporation. In commerce the evaporation is stopped before crystallization begins, and the syrup (" corn-syrup," if maize is the source of the starch) is sold for making candy and for preserving fruits. To understand the action, we must note that, although the STAKCH AND COMPOUNDS MADE THEREFROM 221 formula C 6 HioO 5 represents the composition of starch, its molec- ular weight (the value of which varies) is very large, say (C 6 Hio0 5 )2oo. The water decomposes the large molecules and combines with the material: * -I- xH 2 -* The action is an hydrolysis (Greek, decomposition by water). Glucose is known also as dextrose, and as grape sugar. Brownish crystalline granules found in dried grapes (raisins) are mainly composed of it. When pure, it is almost colorless. It reduces cupric hydroxide, in Fehling's solution, to cuprous oxide. The Sugars. The common sugars are divided into two classes. Two sugars, having the same formula, CeHiaOe, but different properties, are called monosaccharides. Other sugars, having usually twice as many carbon units in the formula Ci 2 H220n, are called disaccharides. The sugars we have occa- sion to mention are the following: Monosaccharides: Glucose (dextrose or grape sugar) CeH^Os. Fructose (fruit sugar) CeH^Oe. Disaccharides: Sucrose (cane-sugar, beet-sugar, saccharose) CizHjaOii. Maltose (formed by action of malt on starch) C^H^Ou. Lactose (milk-sugar, found only in animals) Carbohydrates. Since starch and the sugars are freely changed, one into another, they form one class, the carbohydrates. The word refers to the fact that they usually contain hydrogen and oxygen in the proportions required to form water, and are, therefore, in a sense, hydrates of carbon. Sucrose or Cane- Sugar C^H^On. A number of plants, besides producing starch, are able also to convert carbon dioxide 222 TEXT BOOK OF CHEMISTRY and water into sucrose Ci 2 H 22 On. The sugar-cane and the beet produce exceptionally large amounts of this sugar, which is the one commonly used as table sugar. Maple sugar, obtained by evaporating the sap of the tree, is composed mainly of the same substance. The sugar-cane forms stalks from ten to twelve feet high. The juices are extracted by crushing the plants between rollers. The liquid is evaporated in closed pans. A vacuum maintained in the pans permits the boiling of the solution at a low tempera- ture (about 65 degrees) and prevents the decomposition of a part of the sugar which would otherwise occur. When the syrup cools, the sugar crystallizes and the crystals are freed from the liquid in centrifugal machines. The crystals are brown in color. At the sugar refinery they are dissolved, and the solution is passed through a column of bone charcoal. This absorbs the coloring matter, and the filtrate is once more evaporated and allowed to crystallize. Refined cane-sugar has a faint yellow tint, and a small amount of ultramarine is added to cover up this tint, and give the white appearance which is popularly connected with purity in sugar. The sugar beets, which contain 16 per cent or more of cane- sugar, are sliced and steeped in water to extract the sugar. The liquid contains gummy material in colloidal suspension. This is coagulated and precipitated by adding "milk of lime" (calcium hydroxide Ca(OH) 2 suspended in water) and boiling. Carbon dioxide is then passed through the solution to precipitate the excess of lime: Ca(OH) 2 + C0 2 -> CaC0 3 J + H 2 0. The solution is decolorized with charcoal and evaporated to crystallization in the same way as is the extract from the sugar- cane. Properties of Sucrose. Sucrose crystallizes in four-sided prisms, the form of which is seen in " rock-candy ." It melts STARCH AND COMPOUNDS MADE THEREFROM 223 at 160. It does not reduce Fehling's solution (p. 221). When heated to 200 to 210 it begins to decompose, slowly losing water and leaving a brown, soluble mass called caramel, used in color- ing whisky and soups. When boiled with water, to which a trace of an acid (con- tact agent) has been added, it is hydrolyzed, giving a mixture of the two monosaccharides, glucose and fructose: CuHssOn + H 2 -> CeHiaOe + C^Oe. This mixture of glucose and fructose is called invert sugar and is found in many sweet fruits and in honey. Each sugar interferes with the crystallization of the other, by lowering the freezing-point (p. 114), and so invert sugar is added in making " fondant" candy and candy that is to be "pulled," both of which are intended to remain soft for a longer or shorter tune. With the same object in view, vinegar, lemon juice, or cream of tartar is added to a syrup made from cane-sugar, in order that the acid contained in them may produce some invert sugar and so give a less crystallizable mixture. Prolonged heating has the same effect. Enzymes. Yeast belongs to a low order of plants and consists of minute cells of microscopic size. Its value lies in the fact that, while growing and multiplying, it secretes within each cell small amounts of two very active chemical substances which are dis- solved in the cell contents. These substances are known as zymase and invertase (more strictly, sucrase, meaning an enzyme that splits sugar), and belong to the class of organic materials called enzymes. They produce remarkable chemical changes by their mere presence (contact actions). Alcoholic Fermentation. When a cake of yeast, which contains the living plants, is broken into an aqueous solution of glucose (p. 220), the small amount of zymase causes the gradual decomposition of the sugar. The most favorable temperature is 224 TEXT BOOK OF CHEMISTRY about 30. Bubbles of carbon dioxide soon begin to rise to the surface, and the gas can be led off (Fig. 67) to exhibit its char- acteristic action (p. 201) on limewater. At the same time alcohol C2H 5 OH accumulates in the liquid as the sugar disappears: C 6 H 12 6 -> 2C0 2 t + 2C 2 H 5 OH. The liquid extracted from the yeast cells works as well as does the plant itself, and dead yeast produces the change also. The live plant is employed because then only a few cells are needed to start with, and others rapidly form as the plant grows. Yeast will ferment fructose (CeH^Oe) , with the same result, but more slowly. When placed in invert sugar (p. 223), therefore, it decomposes the glucose first and the fructose afterwards. Zymase does not act upon sucrose (Ci2H 2 20n) . But the invertase (sucrase), which is also contained in the yeast, hydrolyzes the sucrose in the same way as does a dilute acid, giving invert sugar. The latter is then decomposed by the zymase. Hence cane-sugar in solution is decomposed by yeast into alcohol and carbon dioxide, just as is glucose, only more slowly. In the manufacture of wines the glucose contained in the grape juice is fermented by a species of yeast always found on the skins. Common Alcohol C 2 H 5 OH. This alcohol is related to ethane C 2 H 6 containing an hydroxyl (OH) group in place of one unit of hydrogen. Hence, to distinguish it from other alcohols, it is named ethyl alcohol. Ethyl alcohol boils at a lower temperature (78.3) than water. When, therefore, the fermented liquid is distilled, the alcohol FIG. 67. STARCH AND COMPOUNDS MADE THEREFROM 225 comes off first, along with comparatively little water. The stills used in distilleries give 95 per cent alcohol. Ethyl alcohol mixes in all proportions with water. It does not dissolve much cane-sugar or fat, but it dissolves many sub- stances which are little soluble, or are insoluble in water, such as iodine and resins. It is used as a solvent in making varnishes for wood and lacquers for metal, as well as plastics like celluloid, collodion, and artificial silk (p. 384). In Germany it is used for lighting and cooking. Alcohol is not ionized in aqueous solution, and it does not in- teract with dilute acids, bases, or salts. It is, however, more active than the hydrocarbons. Thus, it is easily oxidized, giving, as we shall presently see, acetic acid. With acids, when water is absent, it interacts slowly (see pp. 372, 373). There are other alcohols, of which those most closely related to ethyl alcohol have the general formula C n H 2n +jOH. Commercial Alcohol, Whisky , and Beer. Commercial alcohol is made from the starch of potatoes or maize. Barley, which has been allowed to sprout, and is then dried, is called malt. This contains an enzyme, diastase (more strictly, amylase, mean- ing an enzyme that splits starch), which is able to hydrolyze starch into maltose C^H^On (p. 221). Maltose is further hydrolyzed by yeast to form glucose, and the latter is then decomposed by the zymase into alcohol and carbon dioxide. Whisky is made by treating the starch of rye, maize, or barley in the same way, with subsequent separation of the alcohol (whisky) by distillation. Beer is made similarly from various kinds of grain, particularly barley, except that the fermented liquid is not distilled. Denatured Alcohol. This is alcohol to which 10 per cent of wood spirit and 0.5 per cent of benzine, or 2 per cent of wood spirit and 0.5 per cent of pyridine bases from bone oil, nauseous 226 TEXT BOOK OF CHEMISTRY and poisonous substances, have been added. Pure alcohol, such as can be used in the preparation of spiritous liquors, is subject to a tax of $2.08 per gallon (costs $0.22 and sells at $2.50). De- natured alcohol, since it cannot be so used, is duty-free (in U.S. since 1907). Yet it is equally suitable for varnish-making and other industrial operations, which are thus relieved from the tax with which formerly they were burdened. Acetic Acid HCO 2 CH 3 . This acid is the active substance in vinegar, and has many industrial uses. It is formed by the action of mild oxidizing agents on alcohol. (The active oxidizing agents give with alcohol only carbon dioxide and water.) Vinegar is manufactured by oxidizing alcohol with atmospheric oxygen, using a bacterium (B. Aceti, " mother of vinegar"), or more probably an enzyme which it secretes, as a contact agent. The dilute alcohol, in the form, for example, of "hard" cider (fermented apple juice), is allowed to trickle over shavings in a barrel. The shavings are inoculated with the B. aceti by preliminary wetting with vinegar. Holes in the sides admit a plentiful supply of air (and therefore of oxygen), to the action of which the liquid is exposed by being spread over the surface of the shavings: HOC 2 H 5 + O 2 -* HCO 2 CH 3 + H 2 0. The liquid (vinegar), which issues at the bottom, contains from 5 to 15 per cent of acetic acid, besides coloring and flavoring matters, derived from the fruit juices. Pure acetic acid is prepared by distilling the vinegar repeatedly. Large amounts are obtained also from the liquid distillate, ob- tained by heating wood in the manufacture of charcoal. It is a colorless liquid which boils at 118 and freezes at 16.7. It is a monobasic add. In aqueous solution it is slightly ionized, the radicals being H+ and CO 2 CH 3 -. There are other acids of the same series, the general formula of which is HCO 2 C n H 2n+ i. STARCH AND COMPOUNDS MADE THEREFROM 227 Cellulose (CsHioOs),,. This substance, named cellulose be- cause it forms the walls of the cells, composes much of the frame- work and intricate structure of plants. We are familiar with pure cellulose in the forms of filter paper (see paper manufacture, p. 277) and cotton. The latter consists of fine, hollow tubes of cel- lulose (see Fig. 2, p. 2), large tufts of which surround the seed of the cotton plant. Linen is almost pure cellulose, wood is largely cellulose, and paper pulp is practically all cellulose. Cellulose interacts with very few chemical substances. It is because it thus remains unchanged, by most substances that come in contact with it, that it can be used as a filter paper. When it does undergo chemical change, it acts as if it contained hydroxyl (OH) groups, and behaves therefore chemically like an alcohol (see p. 383). Exercises. 1. What inference do you draw as the composition of tapioca, sago, and ricfe from the facts that they are fruits and when boiled with water and cooled give a jelly-like mass? How should you confirm your inference? 2. (a) Why does a strong solution of sugar boil at a temperature far above that of boiling water? (b) In evaporation why is the boiling-point lower in a vacuum than in air? 3. In fermentation why does not carbon dioxide appear in bubbles at once? 4. What determines the proportion of alcohol in a wine made from pure grape juice? 5. Write the formulae of the acetates of potassium, and of calcium. 6. How do we ascertain that acetic acid in aqueous solution is only slightly ionized (p. 148) ? 7. What products must be formed when paper is burned? Make the equation. CHAPTER XX CHEMICAL EQUILIBRIUM IN spite of its formidable title, this chapter will introduce nothing novel. Its purpose is to collect together and organize more definitely a number of scattered facts and ideas which have already come up in various connections. On this account, however, it will be all the more necessary for the reader to refresh his remembrance of these facts and ideas by re-reading all pages to which reference is made. Reversible Actions. In discussing Deacon's process (p. 134), it was stated that the action comes to rest although a large amount of both of the interacting substances (20 per cent at 345) still remains available: (20 per cent) 4HC1 + O 2 <=* 2H 2 O + 2C1 2 (80 per cent) . Now the materials thus left unused are presumably no less capable of interacting than were the parts which have already reacted. The solution of this mystery lies in the fact (p. 134) that the products themselves interact to reproduce the initial substances (read the equation backwards). Thus two changes, one of which undoes the work of the other, are going on simultaneously. In consequence of this, neither action can reach completion. As we should expect, experiment shows that it makes no difference whether we start with pure chlorine and steam, or with hydrogen chloride and oxygen; the proportions of the four substances found in the tube, after it has been kept at 345 for a sufficient time, are in both cases the same. A general statement may be founded on facts like this, to the effect that a chemical action must remain more or less incomplete when the reverse action also takes place under the same conditions. Two arrows pointing 228 CHEMICAL EQUILIBRIUM 229 in opposite directions are used in equations representing reversible changes.* The foregoing example of a reversible action, and the following examples which very closely resemble it should now be looked up and studied attentively. The discussion in this and the following sections, for which they furnish the basis, cannot otherwise be understood: (1) the interaction of chlorine and water (p. 139), which was fully discussed at the time; (2) the behavior of water vapor (p. 70), of phosphorus vapor, P 4 2P 2 , of sulphur vapor (p. 266), and of iodine vapor (p. 328), giving 2 I. When the action is one which is reversible, but, under the cir- cumstances being discussed, proceeds farther towards completion in one direction than in the other, the arrow will be modified to indicate this fact : C1 2 + H 2 <=> HC1 + HC10 (p. 139). When this relative completeness is due to precipitation or vola- tilization, the fact may be indicated by vertical arrows: NaCl + H 2 S0 4 *=> NaHS0 4 + HC1 f (p. 120). Nad 1 + H 2 S0 4 ^NaHS0 4 + HC1 (p. 121). Actions Which Proceed to Completion. All chemical actions do not belong to the reversible, incomplete class. Many proceed uninterruptedly to exhaustion of one, or all, of the in- gredients. For example, equivalent amounts of magnesium end oxygen combine completely, 2Mg -f O 2 > 2MgO. Here, how- ever, the product is not decomposed even at the white heat pro- duced by the vigor of the union. Indeed, magnesium oxide cannot be decomposed, and the action reversed, at any temperature we can command. The other complete actions, like the decomposi- tion of potassium chlorate (p. 30), are so because they are likewise irreversible. * The reader must avoid the idea that a reversible action is one which goes to completion, and then runs back to a certain extent. This conception would be contrary to the fact, and inexplicable by the kinetic method. 230 TEXT BOOK OF CHEMISTRY Explanation in Terms of Molecules. Restating these facts in terms of the molecules will enable us to reason more clearly about this variety of chemical change. Suppose we start with the materials represented on one side only of such an equation, say the hydrogen chloride and oxygen in that on p. 228. The molecules of these materials will encounter one another frequently in the course of their movements. In a certain proportion of these collisions the chemical change will take place. In the earliest stages there will be few of the new kind of molecules (say, of chlorine and steam), but, as the action goes on, these will increase in number. There will be two consequences of this. In the first place, the parent materials (in this case, hydrogen chloride and oxygen) will diminish in amount, the collisions between their molecules will become fewer, and the speed of the forward action will therefore become less and less. In the second place, the in- crease in the number of molecules of the products will result in more frequent collisions between them, in more frequent occurrence of the chemical change which they can undergo, and thus in an increase in the speed of the reverse action. The forward action begins at its maximum and decreases in speed progressively; the reverse action begins at zero and increases in speed. Finally the two speeds must become equal, and at that point perceptible change in the condition of the whole must cease. The most immediate inference from this mode of viewing the matter is, that the apparent halt in the progress of the action does not indicate any cessation of either chemical change. Both changes must go on, in consequence of the continued encounters of the proper molecules. But since the two changes proceed with equal speeds they produce no alteration in the mass as a whole. In fact, the final state is one of equilibrium, and not of rest, one of balanced activity and not of repose. Hence, chemical changes which are reversible lead to that condition of seemingly suspended action which we speak of as chemical equilibrium. CHEMICAL EQUILIBRIUM 231 Chemical Equilibrium and its Characteristics. The de- tailed discussion of the relations of liquid and vapor (pp. 62, 63- 65), and of saturated solution and undissolved solid (pp. 109, 111- 112), has already familiarized us with the term equilibrium and its significance. We can, in fact, apply to the discussion of any kind of reversible phenomena, the sets of ideas in regard to exchanges of molecules there elaborated. In particular, the reader will note that the three characteristics of a state of equilibrium, developed and illustrated in the case of the physical equilibrium between a liquid and its vapor (p. 62), apply also to a typical case of chemical equilibrium, such as that in Deacon's process now before us. Thus: 1. There are the two opposing tendencies, which ultimately balance one another. Here they are the tendency of the steam and chlorine to produce hydrogen chloride and oxygen, and the tendency of the hydrogen chloride and oxygen to reproduce steam and chlorine by this interaction. 2. At equilibrium the two opposing tendencies or activities are still in full operation, although their effects then neutralize one another. 3. (and this is the chief mark of chemical, as it is of physical equilibrium). The system is in a sensitive state, so that a change in the conditions (temperature and pressure or concentration), even if slight, produces a corresponding change in the state of the system, and does this by favoring or disfavoring one of the two opposing tendencies or apparent activities. Such a change is called a dis- placement of the equilibrium, for the system settles down in a new state of equilibrium with new proportions of the two sets of sub- stances, corresponding to the changed conditions. Thus, in the present instance, a change from 345, where there is 80 per cent of the material in the form of steam and chlorine, to 384 results in the diminution of this proportion to 75 per cent. The equilib- rium is affected by changes in concentration also, as we shall presently see (pp. 232, 234). Now, the foregoing facts show that the key to understanding 232 TEXT BOOK OF CHEMISTRY chemical activities, their magnitudes, their changes, and especially their practical results, must lie in knowing how changes in the conditions affect them. Hence, to the chemist, familiarity with the influence of conditions on chemical phenomena must be of the greatest practical importance. We therefore address ourselves to the discussion of this subject. The "conditions" to be considered are familiar, temperature, and concentration or, in the case of a gas, partial pressure. The " activity" of an action is accurately measured by the speed with which the action proceeds. Thus, if the foregoing section be re- examined, it will be seen that we spoke throughout of the speedy rather than of the tendency or activity. Finally, temperature and other conditions influence also the activities in, and therefore the speeds of, those actions which pro- ceed to completion, and are not reversible. Hence, unless our statements are expressly restricted to reversible actions and to states of equilibrium, they apply to all chemical changes. The Influence of Concentration. In the first place, let us assume that the temperature is constant, and let us confine our attention for the present to the influence of concentration upon a chemical reaction. We have seen (p. 230) that the speed of a chemical change is determined by the frequency with which the molecules of the interacting substances encounter one another. The frequency of the encounters amongst a given set of molecules, resulting in a definite chemical change, will in turn evidently depend entirely upon the degree to which the molecules are con- centrated in each other's neighborhood. Larger amounts of one of the materials, for example, will not result in more rapid chemical action, if the larger amount of material is also scattered through a larger space. Chemical changes, therefore, are not accelerated by increasing the mere quantity of any ingredient, but only by in- creasing the concentration of its molecules. Thus, a large amount of hydrochloric acid with a piece of zinc will generate hydrogen no CHEMICAL EQUILIBRIUM 233 faster than a smaller amount. But substitution of more concen- trated acid will instantly increase the speed of the action. In the second case, the number of molecules of the acid reaching the zinc per second is greater, and this action, being non-reversible, pro- ceeds more rapidly to complete consumption of the zinc. So also, iron burns faster in oxygen (100 per cent) than in air (20 per cent oxygen). With a reversible action the effect on the speed is the same, ex- cepting that the continued activity of the reverse action prevents the direct one from reaching completion. Thus, if, in the action of hydrogen chloride upon oxygen, we introduce into the same space an extra amount of oxygen, this facilitates the formation of steam and chlorine by increasing the possibilities of encounter between molecules of hydrogen chloride and oxygen. At the same time it does not affect (cf. p. 228) the number of encounters in a given time of steam and chlorine mole- cules with one another which result in the reverse transformation. The proportion of chlorine (and steam) formed, therefore, from a given amount of hydrogen chloride will be greater, although the total possible (by complete consumption of the materials) has not been altered, since the quantity of one ingredient only has been increased. The introduction of an excess of hydrogen chloride would have had precisely the same effect. An Experimental Illustration. A reaction in which the effects of different concentrations were carefully studied by Glad- stone (1855) affords a good illustration. If ferric chloride and ammonium or potassium thiocyanate are mixed in aqueous solu- tion, a liquid containing the soluble, blood-red ferric thiocyanate is produced. The compound radicals are (NILi) and (CNS), and the action is a simple double decomposition: FeCla + 3 NEUCNS ? Fe(CNS) 3 + 3 N^Cl. The action is a reversible one, and the mixture is homogeneous, i.e., there is no precipitation. Now, if the two just-named salts are 234 TEXT BOOK OF CHEMISTRY mixed in very dilute solution in the proportions required by the equation, say by adding 20 c.c. of a decinormal solution (p. 169) of each salt to several liters of water, a pale-reddish solution is obtained. When this is divided into four parts, and one is kept for reference, the addition of a little of a concentrated solution of ferric chloride to one jar, and of ammonium thiocyanate to another, will be found to deepen the color by producing more of the ferric thio- cyanate. On the other hand, mixing a few drops of concentrated ammonium chloride solution with the fourth portion will be found to remove the color almost entirely, on account of its influence in accelerating the backward change. The Law of Molecular Concentration. The general prin- ciple discussed and illustrated in this section may be called the law of molecular concentration, and may be stated as follows: In every chemical change the apparent activity, and therefore the speed of the action, is proportional to the molecular concentration of each interacting substance. This holds whether the action is reversible or not. We shall next give a more precise, semi-mathematical formula- tion of this law, as this formulation will be of use later,* and then proceed to illustrate the application of the law, by showing how it explains large classes of actions of which we have already en- countered many examples. ^Formulation of the Law of Molecular Concentration. The mathematical formulation of the law describing the influence of the concentration of the molecules of each participating sub- stance upon the speed of the action is extremely simple. When the actual concentrations of the molecules are specified (in moles, pp. 169, 170, per liter), and the speed is suitably expressed (in moles transformed per minute or per hour), we find that the speed is * This formulation of the law is not required, or referred to in the sections which follow. The section and the following one may therefore be omitted for the present and be taken up in connection with pp. 391 and 406. CHEMICAL EQUILIBRIUM 235 proportional to the concentration of each molecule appearing in the molecular equation for the action. Thus in the interaction of hydrochloric and hypochlorous acids (the reverse of the action of chlorine on water, p. 306), if [HC1] and [HC1O] represent the concentrations of the molecules HC1 and HC1O, and k is a con- stant, and S is the speed, then [HC1] X [HC10] X k = S. Again, for the dissociation of phosphorus pentachloride vapor into phosphorus trichloride and chlorine (p. 241) : PCls PCla + Cl2, if [PC1 5 ] represents the concentration of the PC1 5 molecules, fa is a constant, and Si is the speed of decomposition: Similarly, for the reverse action: PCla + C1 2 -> PC1 5 , if [PC1 3 ] and [C1 2 ] stand for the molecular concentrations of these substances: x [cy x b = &. The constant has a different value in each separate action. It includes the value of the intrinsic affinity or activity of the sub- stances, and the catalytic effect (p. 236), if any, of the materials present. Formulation of the Condition for Chemical Equilibrium. - The foregoing plan may be used further to formulate the con- dition for chemical equilibrium. As we have seen (p. 230), a characteristic of a system in chemical equilibrium is that the speeds of the forward and reverse actions have become equal. If, then, [PCl 5 ] e qm., [PCl 3 ]eqm., and [Cl 2 ]eqm. now represent the molecular con- centrations when the system has reached equilibrium, then, since the speeds are equal: [PCl 3 ]eqm. X [Cl 2 ]eqm. X fa = [PCl 5 ] e qm. X fa, [PCl 3 ] e q m . X [Cl 2 ] eqm . = fa [PCl 5 ]e 2H 2 O + C1 2 ) to run to completion. This might be realized by causing one end of a sealed tube charged with the substances, after the contents had settled down to a condition of equilibrium, to project from the bath in which the whole had been kept at 345 (Fig. 68, which is simply diagram- matic). By cooling this end, a large part of the steam would quickly be condensed in it to the liquid form, while the other sub- stances would remain gaseous. FlG> 68< In other words, the concentra- tion of the water vapor would be greatly reduced. In fact, only the trace of vapor which cold water gives would then be avail- able to interact with the chlorine, and reproduce hydrogen chlo- ride. Meanwhile the decomposition of the latter would go on, and thus, eventually, almost all the water would be found in one end of the tube, and the chlorine, all free, would occupy the rest. By this purely mechanical adjustment the chemical change would therefore be carried from 80 per cent completion to almost absolute com- pletion: 4HC1 + 2 * 2C1 2 + 2H 2 (vapor) <= 2H 2 (liq.) If, on the other hand, arrangements were made to have pow- dered marble, in a sealed bulb of thin glass, enclosed in the tube, we might imagine the very opposite of the above effect to be pro- duced. The breaking of the bulb of marble, when equilibrium had been reached, would provide means for the removal of all the CHEMICAL EQUILIBRIUM 239 hydrogen chloride,* while the other three substances would still be gaseous. Thus, the compound (HC1) having been reduced in concentration to the point of being removed entirely, there would be no direct action to undo the work of the reverse action. The whole chlorine would, therefore, soon have passed through the form HC1. Hence, by another mechanical arrangement, an action which ordinarily could progress to only 20 per cent would be turned into a complete one: 2C1 2 + 2H 2 O <= 2H 2 O + 4HC1 (+CaC0 3 - CaCl 2 + H 2 + C0 2 ). Reversibility Usually Avoided. In every-day chemical work, since our object is usually to prepare some one substance, chemists either avoid chemical changes which are notably re- versible, or adjust the conditions, as is done in the foregoing illustrations, so that the reverse of the action which they desire is prevented. In consequence of this, when carrying out the direc- tions for making familiar preparations, the fact that such actions are reversible at all very readily escapes our notice. Arranging the conditions so that the separation of a solid body by precipita- tion, or the liberation of a gas, takes place, are the two commonest ways of rendering a reversible action complete. Excellent ex- amples of both of these are furnished by the chemical change used in producing hydrogen chloride by the interaction of salt and sulphuric acid, the full discussion of which (p. 120) should now be studied attentively in the light of these explanations (p. 121, foot- note). History. The conceptions discussed in this chapter are not new, although they have come into general use rather recently. The law of reaction speed, and the influence of the concentrations * The hydrogen chloride would be destroyed by interaction with the marble : 2HC1 + CaC0 3 -+ CaCl 2 + CO 2 + H 2 O. The calcium chloride is a solid. The gas, carbon dioxide, does not interact with the other substances, and would not, therefore, interfere with the forma- tion of fresh hydrogen chloride. 240 TEXT BOOK OF CHEMISTRY of the reacting substance thereon (p. 228), was set forth and formulated by Wilhemy as early as 1850. Gladstone (1855) studied quantitatively the influence of concentration in cases of chemical equilibrium (p. 233). The kinetic explanation (p. 230) was developed by Williamson (1851). Finally, the laws of chemi- cal equilibrium were formulated more explicitly and applied more thoroughly by two Swedish chemists, Guldburg and Waage (1864-9). The Influence of Temperature on the Speed of any Re- action. The activity of chemical change, and therefore the speed of all chemical changes, is increased by raising the temperature and diminished by lowering it (cf. p. 30). Thus, zinc displaces hydrogen more rapidly from hot than from cold hydrochloric acid. Different actions are effected in different degrees, and no simple rule accurately defining the effect can be given. Roughly speak- ing, however, a rise of 10 doubles the speed of every action. A rise of 100 will therefore make the speed roughly 1024 times greater. Hence, when the chemist finds that two substances show no evidence of interaction, he infers that there must be either slow action or none, and he seeks to settle the question quickly by heating the mixture. The Influence of Temperature on a System in Equilib- rium. In a reversible change the two opposing reactions are different actions and their speeds are therefore affected in different degrees by the same alternation in temperature. Hence, when the temperature is changed, the relative amount of the two sets of materials present is altered and the equilibrium is displaced. Thus, in Deacon's process, a rise of 40 in the temperature displaces the equilibrium backwards (p. 331), and diminishes the yield of chlo- rine by 5 per cent. In the vapor of phosphorus pentachloride (p. 335), the displacement is in the opposite direction. The vapor is a mixture of the pentachloride with the trichloride and free CHEMICAL EQUILIBRIUM 241 chlorine: PC1 5 <=* PCla + C1 2 . At 200, 51.5 per cent of the material is present as pentachloride and 48.5 per cent as trichloride and chlorine. Raising the temperature to 250 changes the pro- portions to 20 per cent and 80 per cent, respectively. At 300 only 3 per cent of the pentachloride remains. Evidently, here, raising the temperature favors the decomposition of the pentachloride, and therefore increases the speed of its dissociation more than it does the speed of the reunion of the trichloride and chlorine. Fan't HafTs Late. One use of a law is to enable us to answer a question, when we have not in memory the fact constituting the answer, and even when we have never read or heard the fact. The law or rule enables a little reasoning to take the place of a vast amount of memorizing. Thus, to answer the question: Does sodium chloride always have the same composition, it is not necessary to have read and to remember all, or any of the numerous investigations of this substance that have been made. We simply refer the question, mentally, to the law of definite proportions, and say "yes." Now the facts mentioned above are connected by a law which will answer many practical questions in chemistry. When phosphorus trichloride and chlorine combine (to form PC1 5 ), heat is given out. Conversely, when phosphorus penta- chloride dissociates, heat is absorbed: PC1 5 + 30,000 cal. * PCls + C1 2 . Now, when the temperature is raised, the action proceeds in the direction of decomposing more of the pentachloride. That is, the equilibrium is displaced in the direction which absorbs heat. In Deacon's process, we find that the interaction of hydrogen chloride and oxygen liberates heat. 4HC1 + 2 <= 2H 2 + 2C1 2 + 28,000 cal. and in this action raising the temperature drives the equilibrium backwards, and a lowering in the temperature is required to increase the yield of chlorine. 242 TEXT BOOK OF CHEMISTRY The rule is obvious, and applies to all reversible reactions: When the temperature of a system in equilibrium is raised, the equi- librium point is displaced in the direction which absorbs heat. In other words, a rise in temperature favors the interaction of that one of the two sets of materials to which the heat is added (+ sign) in the equation. If the equation happens to be written with a negative heat of reaction (e.g., p. 156), the heat can, of course, be transferred to the other side with its sign changed. This law is known as Van't Hoff's law of mobile equilibrium. This law is of practical value. More than once, in chemical factories, much time and money have been spent on trying to arrange machinery to give a better yield of some substance at a high temperature, when a reference to this law would have shown that the chief change necessary was to use a lower temperature. We shall frequently have occasion to refer to this law. Application to Physical Equilibria. Van't Hoff's law applies also to physical processes. Thus, as the temperature rises, a substance which absorbs heat in dissolving will become more soluble. This is the commoner case, as is shown by the way in which most solubility curves ascend with rising temperature. Conversely, a substance which gives out heat in dissolving is less soluble with rising temperature in a solution already almost saturated with the compound. For example, anhydrous sodium sulphate gives out heat in dissolving, and so its solubility diminishes, with rising temperature. Again, the vaporization of a liquid absorbs heat, and so an in- crease in temperature will increase the pressure, and therefore the concentration of its vapor (p. 62). Le Chatelier's Law. The above mentioned law is really a particular case of a more general one. If some stress (e.g., by change of temperature, pressure, or concentration) is brought to bear on a system in equilibrium, the equilibrium is displaced in the CHEMICAL EQUILIBRIUM 243 direction which tends to undo the effect of the stress. Thus, raising the temperature furthers the change which absorbs heat and therefore would tend to lower the temperature. Increasing the concentration of the molecules pushes the action in the direction which uses up these very molecules (p. 231). Again pressure causes ice to melt, because the water which is formed occupies a smaller volume, and this change tends to relieve the pressure. But pressure will not cause most substances to melt, because usually the liquid form occupies a greater volume and its produc- tion would tend to increase pressure. Summary. In this chapter we have answered three ques- tions : 1. Why do some chemical actions cease, while still incomplete? Answer : They are reversible. 2. What explains the position of the equilibrium point? An- swers: (a) Equal effects of opposed molecular actions; (6) Equality in speed of opposed reactions. 3. What will displace the equilibrium point? Answer: (a) Change in concentration of one (or more) of the substances; (6) Change in the temperature. Exercises. 1. Explain the completeness of the action by which hydrogen chloride and water, respectively, are formed by direct union of the elements. 2. Explain the completeness of the action by which silver chloride (p. 166) is formed. 3. Explain why the decomposition of potassium chlorate is complete. 4. In view of the statement on p. 16, explain why mercuric oxide is completely decomposed by heating. 5. Why can magnetic oxide of iron be reduced completely by a stream of hydrogen (p. 57), and iron oxidized completely by a current of steam (p. 51) ? 244 TEXT BOOK OF CHEMISTRY 6. With the phosphorus pentachloride system, say at 250, what effect would suddenly enlarging the space containing a given amount of the vapor produce? What would be the effect of diminishing the space ? What would be the effect of introducing additional chlorine into the same space (p. 241) ? 7. By what practical means could the degree of dissociation of sulphur vapor (S 8 ) be reduced, without changing the temperature ? 8. What inference should you draw from the fact that: (a) the solubilities of potassium nitrate and of Glauber's salt increase with rise in temperature; (b) that those of calcium hydroxide (p. 110) and triethylamine decrease with rise in temperature? CHAPTER XXI NITROGEN. THE ATMOSPHERE IT is time now to return to the atmosphere, of which the most active component, oxygen, has already been discussed. The other chief component, nitrogen, will lead us to ammonia NH 3 and nitric acid HNO 3 , both of which are of great commercial im- portance, and have interesting derivatives. Occurrence of Nitrogen. Aside from the free nitrogen, which forms nearly four-fifths of the bulk of atmospheric air, much nitrogen is found in nature in combination. Potassium nitrate KN0 3 is formed in the soil by the action of bacteria upon animal matter, and sodium nitrate NaNOs is obtained from an immense deposit in Peru and Chili. Nitrogen is an essential constitu- ent of an important class of organic substances called the proteins, which are found in plants, particularly in the fruit, and in the muscles and other tissues of the animal body. Preparation. Nitrogen may be obtained from the air by simply removing the oxygen. This nitrogen is not pure, however, as it retains about one per cent of other gases the "inert gases" of the atmosphere. The oxygen can be removed by allowing pieces of moist phosphorus (Fig. 69) slowly to oxidize in an enclosed specimen of air. The phosphoric acid H 3 PO 4 and other products of the oxidation of the phosphorus dissolve in the water. Pure nitrogen can be obtained from pure compounds of nitrogen. Thus, ammonia gas may be passed over heated cupric oxide 245 FIG. 69. 246 TEXT BOOK OF CHEMISTRY (Fig. 70), and the water removed by bubbling the gas through sul- phuric acid. Skeleton: NH 3 + CuO - Cu + H 2 O + N 2 . Balanced: 2NH 3 + 3CuO -> 3Cu + 3H 2 O + N 2 . A steady stream of nitrogen is most easily made by heating sodium nitrite and ammonium chloride very gently along with a little water in a flask: NaN0 2 + NH 4 C1 <= NaCl + NH 4 NO 2 -> 2H 2 + N 2 T . The double decomposition is reversible, and the first action might be expected to be only partially com- pleted. But the ammonium nitrite NH 4 N0 2 is unstable, and decomposes as fast as it is formed, so that one of the substances required to re- verse the first reaction is removed, and the reversing action does not occur. Physical Properties. Nitrogen is a colorless, tasteless, and odorless gas. Its density is indicated in the formula N 2 (mol. wt. 2 X 14 = 28). It is very little soluble in water, and is one of the gases which Faraday failed to liquefy. When liquefied it boils at 194. It can also be solidified. Chemical Properties. Nitrogen is chemically a rather indifferent gas. It unites easily with a very few elements, notably some of the most active metals, such as calcium and magnesium. When magnesium burns in the air, the white powder which is formed contains some of the nitride of magnesium Mg 3 N 2 , along with much of the oxide: 2Mg + 2 -> 2MgO. 3Mg + N 2 -> Mg 3 N 2 . The presence of the nitride may be shown by the odor of am- monia, given off when the ash is moistened with water: Mg 3 N 2 + 6H 2 O -> 3Mg(OH) 2 + 2NH 3 T THE ATMOSPHERE 247 The compounds with oxygen, such as NO and HNO 3 , and with hydrogen such as NH 3 , are of immense commercial value, but, not being very stable, they are formed only in traces by direct union of the elements. The processes for utilizing these tendencies to union, feeble as they are, for manufacturing purposes, will be described under the compounds themselves. THE ATMOSPHERE The components of the air may be conveniently divided into regular components and accidental components. The regular components, again, consist of three which are present in practically the same proportions in all samples, and three (namely water, carbon dioxide and dust) which vary markedly in quantity. Components Present in Constant Proportions. The compounds whose proportions are practically invariable are nitrogen, oxygen, and the group of inert gases. When the variable components are removed, the proportions of the constant ones are as follows: By Volume. By Weight. Nitrogen 78.06 75 5 Oxvcen. . 21.00 23.2 Argon 94 1 3 The inert gases, excepting argon, are present in traces only. The Water Vapor. The proportion of water vapor in the air is exceedingly variable. When air becomes cool, the moisture separates in cloud and fog, which are composed of minute drops of liquid water. When much moisture is condensed, the drops are larger and fall as rain. When they fall through a cold region, they freeze to hail. When condensation takes place in air already below 0, the fog is composed of solid, and not of liquid particles. 248 TEXT BOOK OF CHEMISTRY The hexagonal crystalline structures of ice which are deposited form snow. On the other hand, when the weather becomes warm, evapo- ration goes on rapidly, especially in the neighborhood of seas, lakes, or moist country, and the proportion of water vapor in the air may be considerably increased. Humidity. The moisture is usually defined in terms of relative humidity, the standard being the quantity required to saturate the air at the existing temperature. A space filled with] air can take up aqueous vapor only until the partial pressure of ; water vapor becomes equal to the vapor pressure of water (p. 62) at the same temperature. The humidity is then said to be 100 per cent. If the partial pressure actually reached is only half as great as the vapor pressure of water at the same temperature, the humidity is 50 per cent. The average humidity may be placed very roughly at about 66 per cent. At 18 (64.4 F.), the vapor pressure of water is 15.4 mm. (Appendix IV). If the total pressure of the atmosphere were 760 mm., then the air would be saturated with moisture at 18, 15 4 and have a humidity of 100 per cent, when - or about 2 per I t)U cent of it by volume was water vapor. Upon cooling to 0, at which temperature the vapor pressure of water is 4.6 mm., this 4 6 air would retain only-^-, or about 0.6 per cent of moisture. At 18 there would be 16.3 grams of water in a cubic meter of air and at only 4.9 grams. The difference, 11.4 grams (11.4 c.c.), would be precipitated as fog or rain from each cubic meter. Test for Moisture in Air. The presence of moisture in air may be shown by placing any deliquescent (p. 60) salt, such as calcium chloride, in an open vessel. The quantity can be measured THE ATMOSPHEKE 249 by driving a known volume of air slowly through a weighed tube containing dry calcium chloride. It may be ascertained also by noting the temperature to which the air has to be cooled before it becomes saturated and deposits fog or dew. For example, if air at 18 has to be cooled to 11 before it deposits dew, it contains 9.8 mm. of water vapor at a pressure of 9.8 mm. If saturated at 18, it would have contained 15.4 mm. Aq. Its relative humidity was therefore 9.8/15.4, or 63.6 per cent. Moisture and Comfort. The chemical changes occurring in our bodies, and particularly the oxidation of waste and of digested food by oxygen carried by the blood, are accompanied by liberation of heat. Yet our bodies must remain at 9^f F. (37 C.). A rise of a few tenths of a degree (F.) produces notice- able discomfort. Much of the heat is lost by radiation from the surface. The extent of this loss depends upon the surface, which is invariable, and upon the surrounding temperature, which we can not always control. Non-conducting clothes reduce the radiation, and are increased in thickness in cold weather. The real adjustment, however, is accomplished, independently of radiation, by evaporation of water at the surface of the skin. The evaporation of 1 gram of water requires about 540 calories of heat. Evaporation of a single half-ounce (14 g.) of water will therefore lower the temperature of 76| kilograms (168 pounds) of water (or of flesh, which is largely water) by one-tenth of a degree C. (nearly 0.2 F.). Our comfort, then, depends upon the possibility of continual, moderate evaporation from the surface of our bodies. "Much" moisture in the ah* means, to us, not necessarily a great absolute amount, but a near approach to the maximum possible at the existing temperature. So the ratio of the amount present, to the maximum the humidity is the significant fact for a practical purpose, such as feeling comfortable (or drying the wash quickly). 250 TEXT BOOK OF CHEMISTRY Ventilation. In winter, cold, and therefore rather dry air is brought into our rooms. When this air has been heated, its relative humidity is too low, discomfort is felt because there is too much evaporation, and moisture has to be added artificially. Here the moisture afforded by evaporation from our bodies has little effect on the air. In summer, however, the outside air is often already nearly saturated at the temperature of the room. At such times the speed of displacement by the ventilating appliances may not be great enough to keep the relative humidity down, and discomfort will arise from the opposite cause. To relieve it, the evaporation may be promoted by electric fans. They do not remove or add any air, but they stir it, and blow away the moist, nearly saturated, layers next to the skin. The chief purposes of ventilation are, therefore, to supply fresh air, to keep it in motion, and to maintain a humidity that is neither too low nor too high. Most authorities believe that, for crowded rooms, the fresh air should be continuously brought in from the outside. The Carbon Dioxide. The breathing of animals, the com- bustion of coal and wood, and the decay (oxidation) of vegetable and animal matter produce carbon dioxide. The same gas issues from volcanoes, and often in great quantities from the soil in regions which are no longer volcanic. The proportion in the air is therefore greatest in crties and in some volcanic regions, and least in the country and over the sea. It varies from 4 parts in 10,000 in the country, to 1 per cent in crowded rooms. Its presence may be proved in any air, and very quickly in the breath itself, by bubbling the air through calcium hydroxide solution (lime-water). Calcium carbonate is precipitated (p. 201). Carbon Dioxide and Respiration. We draw about half a liter of air into our lungs at each breath, or half a cubic meter per hour. In the lungs some of the oxygen is removed, and some carbon dioxide is added. THE ATMOSPHERE 251 Fresh Air, Per Cent. Expired Air, Per Cent. Oxygen 21.00 15.9 Carbon dioxide 0.04 3.7 A candle flame goes out when the proportion of oxygen has fallen to 16.5 per cent. But air will sustain life until the proportion has fallen to about 10 per cent. Nearly all experts are now convinced that the unhealthiness of over-crowded, " stuffy " rooms is not due to the increase in the proportion of carbon dioxide, which is seldom great enough to do any damage. Nor is it due to " poisons " given off by the lungs or skin. In spite of many experiments the presence of such sub- stances has never been proved they are imaginary. The harm is caused by the stillness of the air, which, as we have seen (p. 250), prevents the removal of the water vapor near the skin, and therefore hinders evaporation. Dust in the Air. A beam of sunlight crossing a dark room can be seen by the light reflected from the particles of dust which all air contains. These are chiefly solid bodies, and are composed of salts, limestone, clay, and other rock materials, of soot and other particles of unburnt fuel, of bits of hay or straw, and of fragments of insects and other debris of plants and animals. They also include living particles, such as bacteria, and spores of plants such as moulds. The latter, when they settle upon food, germinate and give rise to putrefaction. Some of the bacteria also produce disease, when they enter the body at a place where the skin has been damaged by a cut or burn. It is instructive to note that natural soil contains about 100,000 micro-organisms, and good, unfiltered river water from 6000 to 20,000, in each cubic centimeter. Ordinary, pure air contains only 4 to 5 micro-organisms per liter. Most of these bacteria 252 TEXT BOOK OF CHEMISTRY come from the drying of soil and the dispersion of the resulting dust. If dust were not present, we should soon notice its absence. There would be no clouds or rain. It appears that moisture will not condense to fog or rain in air which has been filtered, by being drawn through a wide tube containing a long (20 inches or more) plug of cotton, and has so been freed from dust. The particles act as nuclei, round which the liquid grows at the ex- pense of the vapor. In the absence of dust, the condensation would occur directly upon the surfaces of plants, houses, and animals. Thus, in a dustless atmosphere, an open shed or shelter, or an umbrella, would afford no protection whatever against a wetting. The formation of fog from ordinary air, and its non- formation in filtered air are easily shown in a darkened room (Fig. 71). The flask contains water to saturate the air. When the tube leading to the water pump is opened for an instant, the saturated air in the flask expands and is cooled. In such circumstances, ordinary air gives a fog, brilliantly illumi- nated by the beam of light, while filtered air (dustless) gives none. Air a Mixture. The air does not contain in combined con- dition the various substances we have named. Each of the substances in air shows precisely the same properties which it exhibits when free, separate, and pure. This behavior is charac- teristic of a mixture. Thus, the observed density of the air is precisely that which we find by calculation from the known proportions and several densi- FIG. 71. THE INERT GASES 253 ties of the components. The solubility of each gas is observed to be the same as if it were alone present. Again, when liquefied air is allowed to evaporate in a suitable apparatus, the nitrogen, being more volatile, can be separated almost completely from the oxygen. When the oxygen is, in turn, allowed to evaporate, the carbon dioxide and water remain as solids, frozen by the low temperature. Finally, the exact proportions can not be represented by a chemical formula. This shows that the law that, in chemical compounds, the proportions can be represented by multiples of the atomic weights by whole numbers (p. 77), does not apply to air. This is a sure proof that it is not a chemical compound. In spite of the fact that air is a mixture, the composition of the air is remarkably uniform and constant. The uniformity is due to constant mixing by the winds. The steadiness of the com- position from year to year is due to the fact that, although decay and combustion continually remove oxygen and add carbon diox- ide, vegetation as continually consumes the latter and restores the former (p. 205). The mass of carbon dioxide in the whole atmosphere of the planet, about 2450 thousand million tons, is so great that the amounts added and removed by the agencies just mentioned are small by comparison. THE INERT GASES Argon. It had been shown by Lord Rayleigh that a liter of pure nitrogen weighed 1.2505 g., while a liter of atmospheric " nitrogen" weighed 1.2572 g. The natural inference was that the latter contained a little of some heavier gas. In 1894 Ramsay, in consultation with Rayleigh, attempted to separate this gas by passing the "nitrogen" repeatedly over heated magnesium, and so removing the real nitrogen as solid magnesium nitride Mg 3 N 2 . The remaining gas, about 1 per cent of the whole, was named argon (Greek, lazy or inactive), because it would combine with no other element. 254 TEXT BOOK OF CHEMISTRY Argon has a molecular weight of 39.9 (nitrogen only 28), and when liquefied boils at 186 and freezes at 189.5, Helium. An indifferent gas, previously known to be given off when uranium ores were heated in a vacuum, was found by Ramsay (1895) to be neither nitrogen, nor yet argon. By its spectrum it was recognized to be helium (Greek, the sun), a sub- stance shown in 1868 to be present in the sun. Its molecular weight is 4, so that it is only twice as dense as hydrogen. It was the last gas to be liquefied (by Onnes), and the liquid boils at 268.7 (4.3 Abs.). Like argon, it does not enter into chemical combination. Helium is now being used to fill balloons, because it is not combustible. Other Inert Gases. When liquefied argon was allowed to evaporate, the first vapor coming off was found to contain an- other gas, neon (Greek, new. Mol. wt. 20), along with helium. Careful distillation of the remaining liquid gave two other gases krypton (Greek, hidden. Mol. wt. 83) and xenon (Greek, stranger. Mol. wt. 130). The total amount of these four gases, however, was only 1 part in 80, the remaining 79 parts being pure argon. None of these gases form any compounds. Another gas, given off in minute quantities by compounds of radium, and originally known as the " radium emanation," is now found by Ramsay to be a member of this interesting group of elements, and has been named niton (mol. wt. 222.4). This gas is also indifferent. But it has the startling property of " dis- integrating " of its own accord. It disintegrates into electricity, helium, and elements of somewhat lower atomic weight (see radium). Exercises. 1. Classify (p. 246) each of the reactions repre- sented by equations in this chapter. 2. What are the radicals of sodium nitrite, and what their valences? THE INERT GASES 255 3. At 77 F. the air of a room contains water vapor at a partial pressure of 20 mm. What is the percentage of humidity? 4. What weight of water is contained in a cubic meter (1000 liters) of air saturated at 10 C.? 5. What weight of carbon is contained in the total carbon dioxide in the earth's atmosphere? 6. Air at 18 has to be cooled to 14 before it deposits dew or fog. What is the percentage humidity at 18? 7. Why is the air nearest the ground heated (by the sun) to a higher temperature than the upper air? CHAPTER XXII AMMONIA THE interest in ammonia centers largely in the use of liquefied ammonia for refrigeration, in the employment of the gas in making carbonate of soda, and in the value of its compounds as fertilizers. Manufacture. Ammonia is formed when nitrogenous or- ganic matter is heated, in absence of air. It was formerly made by distilling scraps of hoofs, horns, and hides. The solution of the gas thus obtained was called " spirit of hartshorn." The pungent odor of smoldering feathers, leather, or fur is, therefore, partly due to its presence in the escaping vapors. From the proteins of the original plants, coal derives a considerable proportion of nitro- genous matter. Hence, when coal is distilled for the making of coal gas, or, on a far larger scale, for the making of coke, much ammonia can be separated, by washing with water, from the mix- ture of gases produced. The aqueous solution is separated from the tar, neutralized with sulphuric acid, and evaporated to give the salt, ammonium sulphate (NH 4 ) 2 SO4. NH 3 + H 2 O -> NH 4 OH (ammonium hydroxide). 2NH 4 OH + H 2 S0 4 - (NH 4 ) 2 S0 4 + 2H 2 O. Ammonia from Coal and Shale. While most of the coke in Germany (80 per cent in 1910) is made in " by-product" coke ovens in which the ammonia and innumerable other by-prod- ucts are coRected and utilized, in the United States 83 per cent of the coke was made in " beehive" ovens, in which all the vapors are burned, uselessly, on the spot. Since the war, about 75 per cent of coke is made in by-product coke ovens. Yet 256 AMMONIA 257 ammonium sulphate is being used as a fertilizer in rapidly increas- ing amounts and finds a good market. In 1911 half the 135,000 tons used in this way (U. S.) were imported, paying duty, while ammonia capable of yielding 400,000 tons of ammonium sulphate worth $24,000,000 were wasted by the coke-makers. The distillation of coal is the chief source of commercial am- monia. In Scotland, however, oil-bearing shale is distilled for the purpose of extracting the petroleum, and much ammonia, liberated at the same time, is collected. Formerly it was allowed to escape, but, in the absence of a protective tariff, the competition of petro- leum from American and Russian wells compelled economy. Now, the profit on the sale of the ammonium sulphate pays the whole cost of mining and distilling the shale. Synthetic Ammonia. The latest method of manufacturing ammonia is by 'the direct union of hydrogen and nitrogen. As ammonia gas decomposes almost completely when heated at 700, it must be made below this temperature. The union of the gases is exceedingly slow. But, below 500, in presence of a contact agent such as specially prepared iron or uranium (Haber's process) the gases, delivered at 185 to 200 atmospheres pressure, unite to give about 8 per cent of the compound each time they are led over the metal. The product is dissolved out with water. The required hydrogen is obtained by the action of steam on iron (p. 51), or as a by-product from the making of chlorine and sodium hydroxide (p. 160). The nitrogen is obtained by passing air over heated iron. Preparation in the Laboratory. In the laboratory am- monia is most readily made by heating a mixture of a salt of ammonium, such as the chloride (NI^Cl) or sulphate, and slaked lime Ca(OH) 2 . Ca(OH) 2 + 2NH4C1 -> CaCl 2 + 2NH 4 OH - 2NH 3 + 2H 2 O. 258 TEXT BOOK OF CHEMISTRY The ammonium hydroxide, formed by the double decomposition, immediately decomposes. To free the gas from water vapor, it is' passed through a tower filled with quicklime CaO (Fig. 72) : CaO + H 2 0-+Ca(OH) 2 . Sometimes, a stream of the gas is generated by warming com- mercial ammonium hydroxide solution (aqua ammonia), and dry- ing the gas as above. Liquefied ammonia is obtainable in small iron cylinders, and is a convenient source when much of the gas is re- quired. The liberation by hydrolysis of nitrides, already noted (p. 246), is interesting: Mg 3 N 2 + 6H 2 0-*3Mg(OH) 2 + 2NH 3 . Composition of Ammonia. That ammonia contains nitrogen may be shown by passing the gas over cupric oxide heated in a tube, and collecting the nitrogen over water: 2NH 3 + 3 CuO -* 3Cu + 3H 2 -|- N 2 . FIG. 72. The hydrogen may be liberated by drying the ammonia, if necessary, with soda-lime, and leading it through a tube containing heated magnesium ribbon: 2NH 3 + 3Mg - Mg 3 N 2 + 3H 2 . Physical Properties. Ammonia is a colorless gas. It has a soapy taste, and a very pungent odor. Its density, recorded in the formula NH 3 , indicates that it is only about half as heavy, bulk for bulk, as air (14 + 3 = 17 against 28.95). It is easily liquefied, boiling at 38.5, and exerting a pressure of 6 atmos- pheres at 10. The gas is exceedingly soluble in water (1 vol. AMMONIA 259 Aq. : 1300 vol. of NH 3 at 0). A 35 per cent solution is sold as "concentrated ammonia." The extreme solubility in water may be shown by the "fountain" experiment (Fig. 73). The flask is filled by downward displace- ment of air. The long tube is closed by a short rubber tube and a clip at the bottom (not shown). The "dropper" contains water, and is closed at the tip with soft wax. The drop of water, expelled by pinching the "dropper," dissolves at once so much of the gas that the water rushes in, like a fountain, through the longer tube. The Solubility of Gases. It will aid us in recalling the solubilities of the various gases in water if we summarize the facts about those thus far de- scribed. The gases may be divided into three groups: JflG. 73. 1. Slightly soluble: Oxygen, hydrogen, methane, nitrogen, the inert gases. 2. Moderately soluble: Chlorine, carbon dioxide. 3. Very soluble: Hydrogen chloride, ammonia, sulphur dioxide. The gases of the first group are, in a general way, about ten times as soluble in other liquids (oils, etc.) as in water. Chemical Properties. Ammonia, as we have seen, is not very stable, and decomposes rapidly and almost completely above 700. A discharge of sparks from an induction coil has the same effect more gradually, and so a sample of the gas con- fined over mercury in a closed tube may be shown to double in volume when decomposed. Every two molecules give four: 2NH 3 -3H 2 260 TEXT BOOK OF CHEMISTRY The most characteristic property of ammonia is that it com- bines directly with acids, giving ammonium salts: NH 3 (gas) + HC1 (gas) -V NH 4 C1 (solid particles). 2NH 3 + H 2 S0 4 (liq.) -> (NH 4 ) 2 SO 4 (solid mass). It combines also with water at a very low temperature (98) to give ammonium hydroxide, a white solid : As the solid dissociates above 98, a solution of the substance, made by dissolving the gas in a great excess of water, is the only form of ammonium hydroxide convenient for use. Ammonium Hydroxide. This substance, as indicated by the way in which we have written its formula, is a base. The ions are (OH)~, given by all bases, and (NELO+, ammomwra-ion, which is found also in the salts mentioned above. The latter is a com- pound positive radical, playing the part of a metallic element. As a base, ammonium hydroxide, although rather weak (little ionized), turns red litmus blue, possesses the characteristic soapy taste and feeling, and enters into double decomposition with acids, neutralizing them : NK^OH + HC1 -> H 2 + NH 4 C1. The salts, obtained by evaporation, are, of course, identical with those formed by union of ammonia with the same acids. Ammonium hydroxide used to be known as " volatile alkali/' in deference to the fact that it decomposes into its constituents (NH 3 + H 2 O), both of which are volatile, while the other alkalies (NaOH, etc.) are not volatile ("fixed")* This property was utilized hi the laboratory method of making ammonia (p. 257) . The Salts of Ammonium. The salts are ionized in aqueous solution, giving NH 4 as the positive ion: (NH 4 ) 2 S0 4 4 2NH 4 + + SO 4 = . AMMONIA 261 When heated, dry, in a tube, they are decomposed. give ammonia and an acid: +HC1T, + 4H 2 S0 4 1 Most of them If the acid is also volatile at, or below, a red heat, like sulphuric acid H 2 S04, the whole salt usually vaporizes. These actions are reversible (read the equation backwards). Hence the acid and the ammonia recombine, and the salt condenses again in a cold part of the tube. This behavior helps _ __ us to recognize a salt of ammonium, for the salts of mercury are the only others which behave in this way. Uses of Ammonia. Some of the uses have already been mentioned. The ammonia process for making carbonate of soda is described under the latter sub- stance. The making of ammonium sul- phate from ammonia is sufficiently ex- plained in the preceding sections, and the value of the salt as a fertilizer has been mentioned. Refrigeration by liquid ammonia de- pends upon the fact that liquid ammonia, like any other liquid, absorbs heat in evaporating. It absorbs 260 cal. per gram. To freeze one gram of water at 0, 80 calories have to be subtracted. Thus, 1 gram of liquid ammonia in evaporating will convert about 3 grams of water to ice. The same principle is also largely used for cooling air (in storage rooms for meat, etc.). The machinery is represented diagrammatically in Fig. 74. The ammonia, first admitted from a bomb of liquefied ammonia, is driven by the pump F along the tube E and condenses to liquid form in the tube coiled in the tank AB. Cold water circulates FIG. 74. 262 TEXT BOOK OF CHEMISTRY through AB, and removes the heat produced by the compression and liquefaction of the gas. The liquid ammonia is allowed to drip through the stopcock G into the lower coil. This is kept exhausted by the compresser F, and the liquid ammonia evapo- rates. In doing this, it takes heat from a 30 per cent solution of calcium chloride in water, which does not freeze even at 12. This cooled brine leaves the tank at D, circulates through another tank (not shown) in which the water-filled ice moulds are sus- pended, and returns to C. When used for cooling refrigerating chambers, the brine passes through a system of pipes suitably placed in the cold-room. The whole machinery is made of iron, as copper and brass are corroded by the ammonia. Ammonium hydroxide solution is sold under the name of household ammonia, and is used, in washing and cleaning, to soften the water. Exercises. 1. Classify the reactions shown by equations in this chapter. 2. How could you recognize the nitrogen and the hydrogen (p. 259)? 3. Why not dry ammonia gas with concentrated sulphuric acid or with phosphorus pentoxide? 4. How could you separate a mixture of oxygen and ammonia ? 5. In what relative volumes do ammonia and hydrogen chloride unite? 6. Write in ionic form the equations for the interaction of am- monium hydroxide: (a) with sulphuric acid; (b) with hydro- chloric acid (p. 260) ? CHAPTER xxnr SULPHUR AND HYDROGEN SULPHIDE SULPHUR, the compounds of which have been so often men- tioned, provides us, in sulphuric acid, with a substance which has more extensive and more important applications in commerce than any other chemical. The element sulphur, itself, enters, with potassium nitrate and charcoal, into gunpowder. Vulcanite is a compound of caoutchouc (rubber) and sulphur. Sulphur is employed to destroy fungi on grape-plants, and furnishes sulphur dioxide for bleaching and disinfecting. Sources. The greater part of the sulphur of commerce comes from Sicily, Louisiana and Texas. In Sicily, free sulphur is mixed with pumice and other rocks. When the lumps of rock, obtained by mining or quarrying, are heated by setting fire to the sulphur (there is no coal in Italy), the sulphur melts and runs to the bottom of the kiln. This product is far from pure, and is dis- tilled from iron retorts. The vapor is condensed in chambers of brick, and the liquid is run into moulds, giving roll sulphur. The first vapor condensed, while the chambers are cold, yields flowers of sulphur. In Louisiana the sulphur occurs in a deposit over half a mile in diameter, below 900 feet of clay, quicksand, and rock. It is obtained by means of borings, which permit four pipes, one within the other, to reach the deposit. Water, previously heated under pressure to a temperature of 170, is pumped down the two outside pipes (6 and 8 inches in diameter). After time has been allowed for the melting of a quantity of the sulphur (it melts at 114.5), compressed air is pumped down through an inner, one-inch pipe. The melted sulphur, alone, has twice the specific gravity of the 263 264 TEXT BOOK OF CHEMISTRY water in the outer pipes. But the mixture of air and sulphur has about the same specific gravity, and so flows freely up a three-inch pipe which surrounds the air pipe. The sulphur runs into wooden enclosures, measuring 150 by 250 feet, in which it quickly solidifies. The product is so pure that, for most purposes, no other treatment is required. The output, at Sulphur, Louisiana 500 tons a day from each well and, in all, 250,000 tons annually supplies the whole demand of the United States, and could easily be increased. A number of sulphates, such as gypsum (CaS0 4 , 2H 2 0) and barite (BaS0 4 ), and several sulphides, such as galena (PbS), zinc blende (ZnS), and pyrite (FeS 2 ), are found in large quantities as minerals. The last two sulphides are used in the manufacture of sulphuric acid. Allotropic Forms of Sulphur. Sulphur appears in two different liquid forms, and in two familiar and perfectly distinct solid varieties. The two latter are called, from their crystalline forms, rhombic and monoclinic sulphur. Physical Properties of Rhombic Sulphur. This form is yellow, with specific gravity 2.06. Natural sulphur, roll sulphur, and practically all of most specimens of flowers of sulphur are of this variety, and are identical in all physical properties. Speci- mens of natural sulphur often show the rhombic crystalline form very clearly. All the forms of sulphur are insoluble in water, and all the crystalline forms are soluble in carbon disulphide. Good rhombic crystals are obtained from the solution (Fig. 10, p. 14). The rhombic form is stable when not heated above 95.5. If kept above this temperature, it slowly changes into monoclinic sulphur. Monoclinic Sulphur. This form is obtained most quickly by first melting some sulphur (m.-p. 114.5), and then allowing it slowly to cool. As the temperature is now above 96.5, the SULPHUR AND HYDROGEN SULPHIDE 265 crystals which grow in the liquid are of the monoclinic variety. They are long, transparent, pale-yellow needles (Fig. 75), almost rectangular in section, and bevelled at the points. A brown tint is due to impurities. The specific gravity is 1.96. This form can be kept indefinitely above 96, but, when allowed to cool below that temperature, it slowly becomes opaque, chang- ing into particles of rhombic sulphur. Heat is given out during this change, as it always is in the passage from any state to one stable at a lower temperature. The Two Liquid Forms: Amorphous Sul- phur. When sulphur is melted, and the liquid is heated, two fluid, mutually soluble forms of sulphur are pro- duced. These are known as Sx and S tt or amorphous sulphur. As the temperature rises, the second variety increases in quantity (with absorption of heat) at the expense of the first variety. When the temperature is lowered, the reverse change occurs : Sx <=* S M (amorphous) . If the temperature is lowered slowly, therefore, only mono- clinic sulphur (by crystallization of the Sx) is obtained. But if the liquid is quickly chilled, by pouring into a cold vessel or into cold water, the S M is found as a non-crystalline substance mixed with the crystalline form. The crystalline form can be dissolved out with carbon disulphide, leaving the amorphous sulphur (Greek, without form) which is not soluble. The proportion of S M varies from 3.6 per cent at 120, to 11 per cent at 160 and about 34 per cent at 445 (the boiling-point of sulphur). S M is very viscous, so that, as its quantity increases, the whole mixture becomes thick. The brown color, however, is due to organic impurities. Amorphous sulphur is held to be a super-cooled liquid, and not a true solid, for true solids are all crystalline. At room tempera- 266 TEXT BOOK OF CHEMISTRY ture it changes into rhombic sulphur, but so slowly that the transformation even of a small part of it can be detected (by treating with carbon disulphide) only after the lapse of many months. At 100 the change is complete in less than an hour. Elastic sulphur. When the melted sulphur is chilled, the amorphous sulphur does not at once become hard. When the sul- phur has been heated to a high temperature, therefore, the chilled material consists at first of a sticky, transparent, elastic material, called elastic or plastic sulphur. In the course of forty- eight hours, however, this becomes opaque and hard, because of the separation of the crystalline and the hardening of the amor- phous varieties. Melting and Freezing-Points. Amorphous sulphur, like glass and other amorphous substances, softens when heated, but has no sharp melting temperature. The two crystalline forms have different melting-points, rhombic at 112.8 and monoclinic at 119.25. But these are difficult to observe, as the rhombic begins to turn into monoclinic above 96.5, and both tend to give amorphous sulphur. The latter lowers the freezing-point of the liquid and the melting-points of both solids. Hence, the only temperature which is easy to observe is that at which both the solid forms melt when heated very slowly, and that at which the liquid freezes if cooled very slowly, namely 114.5. This is the so-called natural freezing-point of sulphur. Chemical Properties. The vapor density of sulphur indi- cates that the vapor is a mixture of the molecules S 8 , S 6 and 82, the former diminishing and the latter increasing in number as the temperature is raised. All the metals, excepting gold and platinum, combine with sulphur to form sulphides, and in most cases much heat is given out during the union. Sulphur unites with chlorine to give sulphur mono- SULPHUR AND HYDROGEN SULPHIDE 267 chloride 82012, used in vulcanizing rubber, and burns in oxygen to give sulphur dioxide: * S + 2 ->S0 2 . In these compounds the valence of an atomic weight of sulphur appears to be one (in SjjCy or four (in SO 2 ) . These are exceptional values, however, the common valences being two (in H 2 S, ZnS, etc.) and six (in SO 3 , SO 2 C1 2 , etc.). Moist sulphur is slowly oxidized at ordinary temperatures to sulphuric acid: Skeleton: S + H 2 O + O 2 -H 2 S0 4 . Balanced: 2S -f 2H 2 O + 30 2 - 2H 2 SO 4 . In the equations, the simple formula S is used in place of a molecular formula. The latter is needed only when questions about the volume of the vapor are asked, and sulphur is almost always used only in solid or melted form. Then, too, the vapor contains several kinds of molecules, and using Ss or 8$ would introduce large and inconvenient coefficients. HYDROGEN SULPHIDE H 2 S Occurrence. Sulphur is a constituent of albumen, of which, for example, the white of an egg is composed. When decay takes place within the shell, so that air is excluded and the oxidation which accompanies ordinary decay is prevented, the sulphur gives hydrogen sulphide. The latter can be recognized by its odor. Some mineral waters contain a small amount in solution. Preparation. Hydrogen and sulphur combine so slowly that at 310 the completion of the union requires seven days. A trace may be obtained in a few minutes by leading hydrogen over * Traces of sulphur trioxide are found at the same time. They give minute drops of sulphuric acid, which cause a haziness in the gas when it is formed by this action. 268 TEXT BOOK OF CHEMISTRY sulphur, melted in a bulb (Fig. 76). A strip of paper, dipped in lead acetate solution and placed in the wide part of the tube, is darkened by the formation of lead sulphide PbS (black), while acetic acid is also formed : Pb(C0 2 CH 3 ) 2 + H 2 S -> PbS + 2HC0 2 CH 3 . Laboratory Method. The gas is commonly made by double decomposition, using a sulphide to get the S radical, and an acid for the H radical. Ferrous sulphide, made by heating iron filings and sulphur, is the cheapest sulphide, and it interacts easily with hydrochloric acid or sul- phuric acid: FeS + 2HC1 = FeCl 2 + H 2 S f . The action, like all double decompositions, is reversible. But use of an excess of hydro- chloric acid forces it forward, and the escape of the gaseous hydrogen sulphide reduces the backward action almost to zero. The gas can be made in a flask fitted like that in Fig. 25 (p. 52), or in a Kipp's automatic generator (Fig. 77). It can be collected by upward dis- placement. Valence in Double Decomposition. It will be recalled that the atom of iron may be either bivalent or trivalent. In FeS it is bivalent, and it will be noted in the foregoing equation that it becomes ferrous chlo- ride FeCl 2 , in which it is still bivalent. It is a general rule that, in a simple double decomposition, each element and radical carries the same valence through the action. Physical Properties. Hydrogen sulphide is a colorless gas with an odor recalling rotten eggs. It is rather easily liquefied, FIG. 77. SULPHUR AND HYDROGEN SULPHIDE 269 and the liquid boils at about -60 and freezes at -83. The density, implied in the formula H 2 S, shows that 22.4 liters weigh 32 + 2 or 34 g., so that the density is only one-sixth greater than that of air (of which 22.4 1. weigh 28.95 g.). The gas is moderately soluble in water (360 vols. in 100 vols. Aq), a property which enables us to carry out many reactions of the gas upon substances in solution. Physiological Properties. Care must be taken to allow as little of the gas as possible to escape into the air, and all work with it should be done in a well- ventilated hood. The proportion must reach 1 part in 200 of air, however, before fatal results follow breathing the mixture. Chemical Properties. 1. The gas burns in the air, giving water and sulphur dioxide : Skeleton: H 2 S + 2 -> H 2 O + S0 2 . Balanced: 2H 2 S + 3O 2 - 2H 2 O + 2SO 2 . 2. The compound is not very stable. When heated, for example, in the interior of its own flame, it is decomposed into free sulphur and hydrogen. A cold porcelain dish (Fig. 78) placed in the flame will condense some of the sulphur on its surface. 3. On account of its instability, and the ease with which it gives up hydrogen, the gas is a reducing agent. Thus, when jars of hydro- gen sulphide and sulphur dioxide are placed mouth to mouth, a deposit of sulphur gradu- ally appears: S0 2 + 2H 2 S -* 2H 2 + 3S. FlG> 78 ' Part of the free sulphur found in nature seems to be liberated by the action of these gases, both of which are found in volcanic regions. The gases must be moist, for, without water vapor as a contact agent, no interaction occurs. 270 TEXT BOOK OF CHEMISTRY In this action the sulphur dioxide loses its oxygen. We say that the H 2 S was oxidized by the S0 2 , or that the S0 2 was reduced by the H 2 S. As we have already noted, every reduction involves also an oxidation. Chemical Properties An Acid. The aqueous solution is an add, and hence the compound is frequently called hydro- sulphuric acid. It turns faintly tinted litmus paper distinctly pink. The poor conductivity of the solution shows the substance to be little ionized, and therefore a weak acid. Like all acids, it enters into double decomposition with bases and salts. A number of these actions are used in analytical chemistry. Thus, with cupric sulphate solution, we get cupric sulphide (black), and with antimony trichloride antimony tri- sulphide (orange), both as precipitates: CuS0 4 + H 2 S <= CuS i + H 2 SO 4 . 2SbCl 3 + 3H 2 S * Sb 2 S 3 J, + 6HC1. If we keep in mind that the action really takes place between the ions, the first equation assumes the simpler form: The solution surrounding the precipitate, if dilute, contains chiefly the ions H + and SO4 = , along with a few molecules of H 2 SC>4. Metals displace hydrogen from the solution, but, on account of the small amount of hydrogen-ion present, the action is very slow. A striking fact, however, is the displacement of hydrogen from hydrogen sulphide by metals like silver, which do not displace hydrogen from most acids (see activity list, Appendix V). Here, the lack of activity of the silver is made up by the extreme stability and insolubility of the sulphide which is formed: 2Ag + H 2 S-Ag 2 S + H 2 . The tarnishing of silver in the household is due to this action, the hydrogen sulphide coming, probably, from slight leaks in the gas SULPHUR AND HYDROGEN SULPHIDE 271 pipes. Illuminating gas should be free from hydrogen sulphide, but usually contains a trace of it. Sulphides. As we have seen, many sulphides of metals are found as minerals. Since most sulphides are insoluble, many can be made by double decomposition. They may also be prepared by reduction of sulphates. Thus, when sodium sulphate is heated on a piece of charcoal (such as a half -burnt match) the sulphide is formed : 4C -> Na*S + 4CO. It will be observed that in sulphides, H 2 S, Na^S, ZnS, CuS, and so forth, sulphur is invariably bivalent. Carbon Bisulphide 082. This compound is an important solvent for sulphur, caoutchouc (rubber), and other substances which do not dissolve in water. It is manufactured by heating coke and sulphur together to a very high temperature. Electricity is used as the source of the heat, the resistance of the coke causing the latter to glow. The sulphur (vapor) mixed with the coke combines with the latter, and carbon disulphide passes off as vapor and is condensed. The liquid boils at 46, and is highly inflam- mable: Carbon disulphide is used as a solvent. Large quantities are employed also in the destruction of prairie dogs and for freeing grain elevators of rats and mice. Exercises. 1. Write equations for the union of aluminium and of zinc with sulphur. 2. What experiments should you use to recognize a piece of sulphur ? 3. In what proportions by volume do, (a) sulphur dioxide and hydrogen sulphide, (6) oxygen and hydrogen sulphide interact? 272 TEXT BOOK OF CHEMISTRY 4 Make an ionic equation for the precipitation of antimony trisulphide. 5. Would equal weights of rhombic, monoclinic, and amorphous sulphur give out equal or different amounts of heat on burning? If different, which would give the most and which the least? 6. What would be the effect of passing hydrogen sulphide through a red-hot tube? CHAPTER XXIV OXIDES AND OXACIDS OF SULPHUR, PERIODIC SYSTEM THERE are two familiar oxides, namely sulphur dioxide or sul- phurous anhydride SO 2 , and sulphur trioxide or sulphuric anhy- dride S0 3 . Each of these dissolves in water and combines with it to form an acid, just as does carbon dioxide (p. 201). The former gives sulphurous acid H 2 0,S02 or H 2 SO 3 , and the latter sulphuric acid H 2 O,S0 3 or H 2 SO 4 . Acidic and Basic Oxides. As we have seen, an oxide, like carbon dioxide CO 2 (p. 201) or sulphur dioxide S0 2 , which com- bines with water to form an acid, is said to be the anhydride of the acid. The oxides of the non^metallic elements, when they combine with water, in so doing invariably form adds. In the next five chapters we shall meet with other examples (e.g., C1 2 0, N 2 0s, P 2 0s, SiO 2 , etc.). On the other hand, the oxides of metallic elements, when they are able to combine with water, very fre- quently give bases (e.g. Ca(OH) 2 from CaO. For convenience, therefore, we shall often speak of an oxide as an acidic oxide or a basic oxide, as the case may be. Nomenclature. The acids and salts within one group are distinguished by the terminations of, and prefixes to, their names. Thus we have : Hyposulplmrous acid H 2 S 2 4 . Sodium hyposulphite Na 2 S 2 O4. Sulphuroits acid H 2 SO 3 . Sodium sulphate Na 2 SO 3 . Sulphuric acid H 2 S0 4 . Sodium sulphate Na 2 SO 4 . Persulphurc'c acid H 2 S 2 Os. Sodium persulphate Na 2 S 2 Og. 273 274 TEXT BOOK OF CHEMISTRY The proportion of oxygen to the other elements is at the basis of the system. The terminations ous and ite indicate less oxygen than ic and ate. The prefix hypo (Greek, below) implies still less oxygen. The per-acid contains the most oxygen. Knowledge of one group, however, does not enable us to attach names to given formulae from another group. Thus, HC10 2 and HC1O 3 are chlorous and chloric acids, and HC10 and HC10 4 the hypo- and per- acids, respectively. The names of compounds containing only two elements (the binary compounds) end in ide: Zinc sulphide ZnS, magnesium nitride Mg 3 N 2 , calcium carbide CaC 2 , sodium chloride NaCl, and the oxides CaO, etc. SULPHUR DIOXIDE AND SULPHUROUS ACID Preparation of Sulphur Dioxide SO 2 . In commercial prac- tice sulphur dioxide is obtained in three ways : 1. By burning sulphur. The product is used for all purposes to which sulphur dioxide is put. 2. By burning natural sulphides, such as pyrite: Skeleton: FeS 2 + 2 -> Fe 2 O 3 + S0 2 . Balanced: 4FeS 2 + 1 1O 2 -> 2Fe 2 3 + 8S0 2 . With fairly pure pyrite the combustion has only to be started. But with some sulphides, like zinc blende ZnS, which is used as a source of sulphur dioxide as well as of zinc, the air must be strongly heated to maintain the combustion: 2ZnS + 30 2 -> 2ZnO + 2SO 2 . Forced combustion of an ore, like this, is called roasting, or calcining, and is the first stage towards obtaining the metal. The oxide is subsequently reduced by heating with coke. The sulphur dioxide obtained in the process of roasting ores is not always used for any purpose, but, when used, it is made into sulphuric acid. One reason for this is that it is too much diluted SULPHUR DIOXIDE AND SULPHUROUS ACID 275 to be conveniently employed for other purposes. The amount of admixed nitrogen (from the air) alone is enormous. Thus, in the case of zinc sulphide (see equation, above), 3O2 give 2S02, or 3 volumes of oxygen give 2 volumes of the oxide. But each volume of oxygen is accompanied by four volumes of nitrogen, so that, in the end, the 2 volumes of the dioxide are mixed with at least 12 volumes of nitrogen. In practice the dilution is much greater than this. 3. By dropping concentrated sulphuric acid into red-hot iron retorts: 2H 2 S0 4 -> 2H 2 + 2S0 2 + O 2 . Sulphur dioxide made in this way is employed for bleaching, and for the preparation of sulphites (used in paper-making). 4. In the laboratory a steady stream of the gas is easily obtained by dropping hydrochloric acid upon crystals of sodium-hydrogen sulphite (Fig. 25, p. 52): NaHSO 3 + HC1 - NaCl + H 2 SO 3 , (1) - (2) This method takes advantage of the fact that sulphurous acid, like ah 1 acids, can be made by double decomposition (equation 1), and that this acid is unstable and decomposes (equation 2) when there is not a large excess of water present. Physical Properties. The usual six physical properties may be noted: The gas is colorless, but has a characteristic taste and odor. It has a density considerably greater than that of air (SO 2 = 64 against 28.95). It can be liquefied below 165 (the crit. temp.) and the liquid boils at 8. As the pressure required at 20 is only 3J atmospheres, the liquid can be kept in bottles like syphons, or in sealed tin cans. It is extremely soluble in water (about 50 vols. : 1 vol. Aq). The solution is sulphurous acid. 276 TEXT BOOK OF CHEMISTRY Chemical Properties. Sulphur dioxide is very stable. It combines with water giving a solution of sulphurous acid. The gas is used in bleaching straw, silk, and wool. The bleaching action seems largely to consist in combination with the coloring matter, to give a colorless compound. Hence straw hats recover the yellow color of straw by exposure to light, which slowly liberates the sulphur dioxide. It kills fungi and other organisms, but has little effect if too much diluted with air. Formaldehyde is displacing it as a disinfectant. Properties of Sulphurous Acid H 2 SO 3 . The acid is un- stable, and the solution smells strongly of sulphur dioxide. Sulphurous acid, in aqueous solution, shows all the properties of an acid very distinctly; it turns litmus red, gives hydrogen with active metals, and gives double decompositions with bases and salts. Being rather easily convertible into sulphuric acid H 2 S04, sulphurous acid is a reducing agent. Thus oxygen from the air acts slowly upon the solution: 2H 2 S0 3 + 2 ->2H 2 S0 4 and iodine is turned into hydrogen iodide: H 2 S0 3 + H 2 + I 2 - H 2 S0 4 + 2HI. Sulphites and Bisulphites. Sulphites are formed by neutralization : 2NaOH + H 2 S0 3 - Na 2 S0 3 + 2H 2 O. With excess of sulphur dioxide passed into the solutions of the bases, the acid sulphites are formed: NaOH + H 2 S0 3 -> NaHS0 3 + H 2 0. Ca(OH) 2 + 2H 2 SO 3 - Ca(HS0 3 ) 2 + 2H 2 O. Such acid salts are known in commerce as bisulphites, because, the proportion of the metal being half that in a sulphite, the pro- portion of the sulphite radical is, relatively, twice as great. SULPHUR DIOXIDE AND SULPHUROUS ACID 277 Dibasic Acids. Acids containing two atoms of hydrogen in each molecule are called dibasic acids. H 2 CO 3 (p. 199), H 2 S, H 2 S0 3 and H 2 S04 are such acids. Each molecule is able to react with two molecules of a base like sodium hydroxide, as may be seen in the first of the equations in the preceding section. When half the quantity of the base is used, an add salt (p. 276) is produced, as the two other equations show. Phosphoric acid H 3 P04 is a tribasic acid. Hydrochloric and nitric acids HN0 3 are monobasic. Paper Manufacture. Paper is composed of cellulose (CeHioOs)^ and is made from a mixture of cotton or linen pulp and wood pulp the cheapest varieties from the latter alone. The wood is cut into chips and heated (" cooked") with a solution of calcium bisulphite Ca(HSOs) 2 . This dissolves out the lignin, which, together with cellulose, makes up the solid part of its structure. The pulpy material is then washed, beaten with water to reduce it to minute shreds, and bleached with very dilute chlorine-water. The pure cellulose, now paper pulp, suspended in water, is spread on screens, drained, pressed, and dried. During the process other substances are usually added. Thus size (glue or gelatine) prevents the ink from running; pulverized calcium sulphate (gypsum), and other white solids (" loading") give body to the paper and make possible the subsequent production of a smooth surface by rolling (" calendering"). Ultramarine (blue) and other colored powders are added to the pulp when special tints are required. Other Uses of Sulphurous Acid. To prevent the growth of fungi or other organisms, wine casks are fumigated with sulphur dioxide before being filled. Dried peaches and apples are prepared by exposing slices of the fruit on trays to sulphur dioxide. The sulphurous acid produced bleaches the fruit, keeps insects away, and prevents the formation of dark-colored substances during the subsequent drying. In the United States the law (1907) permits 278 TEXT BOOK OF CHEMISTRY the use of natural coloring matters, and of seven harmless coal-tar dyes, in the preparation of preserved fruits and candy for the market. SULPHUR TRIOXIDE AND SULPHURIC ACID Contact Process for Sulphuric Acid. Although the forma- tion of sulphur trioxide SO 3 is accompanied by the liberation of much heat, sulphur dioxide and oxygen, even when heated to- gether, unite very slowly. Ozone, however, combines with the former readily. The interaction of sulphur dioxide and oxygen is hastened by many substances, such as glass, porcelain, ferric oxide and, more especially, finely divided platinum, which remain themselves un- changed and simply act as catalytic agents. The contact process, as this is called, has been rendered available for the commercial manufacture of sulphur trioxide by Knietsch (1901). The chief features of the process are: (1) The complete removal of arsenious oxide, dust, and other impurities derived from the calcining of pyrite or some other mineral sulphide, the minutest traces of which " poison" the catalytic agent and soon render it absolutely in- operative. (2) The preliminary passage of the cold mixture of gases over the outside of the pipes containing the contact agent. This removes part of the heat generated by the action, $Q 2 + O + S0 3 + 22,600 cal., going on inside, and keeps the temperature of the interior at 400. Below 400, the union is too slow; above 400, the reverse action is strengthened (van't Hoff's law, p. 241), and the union is incomplete. At 400, 98-99 per cent of the materials unite; at 700 only 60 per cent, at 900 practically none. Twice the quantity of oxygen theoretically needed is employed. The vaporous product, mainly 1 vol. O 2 to 2 vol. SOs (gas), is condensed by being led into 9799 per cent sulphuric acid, and the concentration of the liquid is constantly maintained at this point by the regulated influx of water. Oleum (q.v.) is also made by omit- SULPHUR TRIOXIDE AND SULPHURIC ACID 279 ting to add water. The trioxide is thus chiefly used for immediate conversion into sulphuric acid. The efficiency of the contact agent depends on the amount of surface it presents to the gases. The action may be illustrated by applying asbestos in a solution of chloroplatinic acid and then heating the mineral in the Bunsen flame: H 2 PtCl 6 - Pt + 2HC1 t + 2C1 2 t The platinum is thus spread in a fine grey powder on the fibers of the asbestos. The latter is placed in a tube (Fig. 46, p. 134), where a mixture of oxygen (or air) and sulphur dioxide may be passed over the heated (400) material. 2SO 2 + 2 -2S0 3 . The sulphur trioxide issues as vapor at the other end of the tube, where its presence is recognized by the dense fumes (droplets of sulphuric acid), produced when it meets the moisture in the air. The vapor can be condensed to liquid form in a cooled flask. It would seem to be simpler to dissolve the gaseous sulphur tri- oxide in water, to give sulphuric acid H 2 O -f- SO 3 > H 2 S04, rather than in 98 per cent sulphuric acid, but this cannot be done. The mixture 2 + 2SO 3 is very incompletely absorbed by water. When a bubble of this mixture enters water, the latter evaporates in the attempt to saturate the space occupied by the bubble with water vapor (p. 62 and Appendix IV). The water, however, combines with the sulphur trioxide to form a fog, consisting of droplets of liquid sulphuric acid, and so more and more water evaporates into the bubbles. The molecules of S0 3 , so long as they remain gaseous, move with great velocity, namely 292 meters per second at room temperature, and still faster in this hot gaseous mixture. Hence, all the molecules that escape combination with the water vapor, strike the wall of the bubble, and combine with the water in a few seconds. The droplets of sulphuric acid, forming the fog, however, are not molecules but large aggregates of molecules. They do not 280 TEXT BOOK OF CHEMISTRY therefore move like the molecules of a gas, but are absolutely stationary. Hence, after the gaseous sulphur trioxide has dis- solved, the droplets of fog, carried by the excess of oxygen, can be bubbled through a whole series of vessels of water in succession without any appreciable number of the droplets being dissolved. The same fog can be shaken in a flask with water, violently and continuously, without any appreciable solution. When the water is thrown, by the shaking, through the oxygen, the oxygen is split up by the water, and driven about, but the fog particles move with the oxygen, so that the water never reaches them. On the other hand, when the mixture of gases bubbles through 98 per cent sulphuric acid, as is done in practice, there is no water available, the sulphur trioxide remains gaseous, and its rapidly moving molecules in a few seconds have all plunged into the sulphuric acid and combined with it, either uniting with the 1 to 3 per cent of water present, or, when oleum is made, uniting with the sulphuric acid to form pyrosulphuric acid : SOa + H 2 S04 H 2 S 2 O 7 . This case affords an admirable illustration of the importance of physics in practical chemistry (p. 21). The chemical reaction occurs with water, but the physical condition of the fog of sulphuric acid prevents its dissolving and, if water were used in a factory, a large proportion of the sulphuric acid would pass with the excess of oxygen into the air and be lost. In fact, it would kill vegetation, and make life unbearable in the neighborhood. Chamber Process for Sulphuric Acid. Although salts of sulphuric acid, such as calcium sulphate CaS0 4 , are exceedingly plentiful in nature, the preparation of the acid by chemical action upon the salts is not practicable. The sulphates, indeed, interact with all acids, but the actions are reversible. The completion of the action by the plan used in making hydrogen chloride (p. 120), involving the removal of the sulphuric acid by distillation, would be difficult on account of the in volatility of this acid. It boils at 330; and suitable acids, less volatile still, which might be used to SULPHUR TRIOXIDE AND SULPHURIC ACID 281 liberate it, do not exist. We are therefore compelled to build up sulphuric acid from its elements. The gases, the interactions of which result in the formation of sulphuric acid, are: water vapor, sulphur dioxide, nitrous anhy- dride N 2 0a * (see index), and oxygen. These are obtained, the first by injection of steam, the second usually by the burning of pyrite, the third from nitric acid HNO 3 , and the fourth by the introduction of air. The gases are thoroughly mixed in large leaden chambers, and the sulphuric acid forms droplets which fall to the floors. In spite of elaborate investigations, instigated by the extensive scale upon which the manufacture is carried on and the immense financial interests involved, some uncertainty still exists in regard to the precise nature of the chemical changes which take place. According to Lunge, supporting the view first sug- gested by Berzelius, the greater part of the product is formed by two successive actions, the first of which yields a complex com- pound that is decomposed by excess of water in the second: 0-H H 2 + 2S0 2 + N 2 3 + 2 -> 2SQ, ( (1) O-NO The group NO, nitrosyl, is found in many compounds. Here, if it were displaced by hydrogen, sulphuric acid would result. Hence this compound is called nitrosylsulphuric acid: 0-H OH +N 2 3 . (2) The equations (1) and (2) are not partial equations for one inter- action, but represent distinct actions which can be carried out separately. In a properly operating plant, indeed, the nitrosyl- * This gas is unstable, breaking up in part into nitric oxide NO and nitro- gen tetroxide NO 2 : N 2 O 3 <=* NO -f- NO 2 . In this process, however, the mix- ture behaves as if it were all N 2 O 3 , and so only nitrous anhydride is named in this connection. 282 TEXT BOOK OF CHEMISTRY sulphuric acid is not observed. But when the supply of water is deficient, white "chamber crystals," consisting of this substance, collect on the walls. The explanation of the success of this seemingly roundabout method of getting sulphuric acid is as follows: The direct union of sulphur dioxide and water to form sulphurous acid is rapid, but the action of free oxygen upon the latter, 2H 2 S0 3 + 2 2H 2 S0 4 , is exceedingly slow. Reaching sulphuric acid by the use of these two changes, although they constitute a direct route to the result, is not feasible in practice. On the other hand, both of the above actions, (1) and (2), happen to be much more speedy, and so, by their use, more rapid production of the desired substance is secured at the expense of a slight complexity. The progress of the first action is marked by the disappearance of the brown nitrous anhydride and, on the introduction of water, the completion of the second stage results in the reproduction of the same substance. The nitrous anhydride takes part a large number of times in these changes, and so facilitates the conversion of a great amount of sulphur dioxide, oxygen, and water into sul- phuric acid, without much diminution of its quantity. Some is lost, however. The loss of nitrous anhydride is made good by the introduction of nitric acid vapor into the chamber. This acid is made from con- centrated sulphuric acid and commercial sodium nitrate NaN0 3 : NaN0 3 + H 2 S0 4 <=* HN0 3 1 + NaHS0 4 . On account of the volatility of the nitric acid, a moderate heat is sufficient to remove it from admixture with the other substances, and its vapor is swept along with the other gases into the apparatus. The initial action which the nitric acid undergoes : H 2 O + 2S0 2 + 2HNO 3 - 2H 2 S0 4 + N 2 3 , may be written, to show the anhydride of nitric acid: H 2 + 2S0 2 + H 2 0,N 2 6 -> 2H 2 S0 4 + N 2 3 . SULPHUR TRIOXIDE AND SULPHURIC ACID 283 The two molecules of water, one actually, the other potentially, present, with the two molecules of sulphur dioxide, can furnish two molecules of sulphurous acid (H 2 SO 3 ) . The NzO$ in passing to the condition ^Oa gives up the two units of oxygen required to con- vert this sulphurous acid into sulphuric acid. Details of the Chamber Process. The sulphur dioxide is produced in a row of furnaces A (Fig. 79) . When good pyrite is used, the ore burns unassisted (p. 274), while impure pyrite and FIG. 79. zinc-blende ZnS have to be heated artificially to maintain the com- bustion. The gases from the various furnaces pass into one long dust-flue, in which they are mingled with the proper proportion of air, and deposit oxides of iron and of arsenic, and other materials which they transport mechanically. From this flue they enter the Glover tower G, in which they acquire the oxides of nitrogen. Having secured all the necessary constituents, excepting water, the gases next enter the first of the lead chambers, large structures lined completely with sheet lead. These measure as much as 284 TEXT BOOK OF CHEMISTRY 100 X 40 X 40 feet, and have a total capacity of 150,000 to 200,000 cubic feet. As the gases drift through these chambers they are thoroughly mixed, and an amount of water considerably in excess of that actually required is injected in the form of steam at various points. The acid, along with the excess of water, condenses and collects upon the floor of the chamber, while the unused gases, chiefly nitrous anhydride and nitrogen, the latter derived 'from the air originally admitted, find an exit into the Gay-Lussac tower L. This is a tower about fifty feet in height, filled with tiles, over which concentrated sulphuric acid continually trickles. The object of this tower, to catch the nitrous anhydride and enable it to be reemployed in the process, is accomplished by a reversal of action (2) above. The acid which accumulates in the vessel at the bottom of this tower contains the nitrosylsulphuric acid, and by means of compressed air is forced through a pipe up to a vessel at the top of the Glover tower G. When this "nitrous vitriol" is mixed with dilute sulphuric acid from a neighboring vessel, by allowing both to flow down into the tower, the nitrous anhydride is once more set free by the interaction of the water in the dilute acid (action (2)). The Glover tower is filled with broken flint or tiles, and the heated gases from the furnace acquire in it their supply of nitrous anhy- dride. Their high temperature causes a considerable concentra- tion of the diluted sulphuric acid as it trickles downward. The acid, after traversing this tower, is sufficiently strong to be used once more for the absorption of nitrous anhydride. To replace the part of the nitrous anhydride which is inevitably lost, fresh nitric acid is furnished by small open vessels N t contain- ing sodium nitrate and sulphuric acid, placed in the flues of the pyrite-burners. About 4 kg. of the nitrate are consumed for every 100 kg. of sulphur. The acid which accumulates upon the floors contains but 60 to 70 per cent of sulphuric acid, and has a specific gravity of 1.5-1.62. The excess of water is needed to facilitate the second action. It is SULPHUR TRIOXIDE AND SULPHURIC ACID 285 required also in order that the acid upon the floor may not after- wards absorb and retain the nitrous anhydride, for this substance combines with an acid containing more than 70 per cent of hydro- gen sulphate. This crude sulphuric acid is applicable directly in some chemical manufactures, such as the preparation of superphosphates (q.v.). Concentration is effected by evaporation in pans lined with lead, which are frequently placed over the pyrite-burners in order to economize fuel. The evaporation in lead is carried on until a specific gravity 1.7, corresponding to 77 per cent concentration, is reached. Up to this point the sulphate of lead formed by the action of the sulphuric acid produces a crust which protects the metal from further action. When a stronger acid is required, the water is driven out by heating the sulphuric acid in vessels of glass or platinum, or even of cast iron. Iron acts upon dilute sulphuric acid, displacing the hydrogen-ion, but not upon concen- trated sulphuric acid, which is not ionized. Commercial sulphuric acid, oil of vitriol, has a specific gravity 1.83-1.84, and contains about 93.5 per cent of hydrogen sulphate. Physical Properties of Sulphuric Acid. The acid is a colorless, oily liquid of sp. gr. 1.84, which freezes to a solid at 10. It mixes with water in all proportions, and much heat (heat of solution) is given out when it dissolves. It boils at 330, but the vapor contains 34 per cent of free water and sulphur trioxide, which recombine when it cools. When heated, it loses sulphur trioxide, but retains the water, and so finally boils at 338. Chemical Properties. 1. The acid is more stable than sul- phurous acid, but decomposes largely at the boiling-point. 2. Sulphuric acid is much more active as an acid than is sul- phurous acid, but inferior in this respect to hydrochloric HC1 and nitric acids HN0 3 . Like other active, soluble acids, its solution turns litmus red, gives hydrogen upon addition of active metals, 286 TEXT BOOK OF CHEMISTRY and enters into double decomposition with bases and salts. Thus, insoluble sulphates are precipitated: BaCl 2 + H 2 S0 4 -> BaS0 4 I + 2HC1. Of course, the same product, barium sulphate, would be obtained with any substance containing the radical SO*: I + 2NaCl and the action is used as a test for this ion. As other salts of barium are insoluble in water, the sulphate is recognized by the fact that it is too insoluble to be acted upon by dilute pure hydro- chloric acid or nitric acid. The other insoluble salts of barium interact with these acids and dissolve. The addition of one of these acids is therefore part of the test for S0 4 -ion. On account of its high boiling-point, the double decompositions of the concentrated acid can be used for preparing more volatile acids: H 2 S0 4 -NaHS0 4 + HCl | (gas at room temp.). H 2 SO 4 -NaHS04 + HN03 t (volatile at 86). 3. Concentrated sulphuric acid combines with water to form a stable hydrate H 2 S0 4 , H 2 0. Hence it removes the elements of water from many substances containing hydrogen and oxygen, and is called a dehydrating agent. Thus, paper (cellulose). moistened with the acid and warmed, turns black from the libera- tion of carbon. Sugar (C^H^On) is decomposed even more easily: C 6 H 10 5 -6C + 5H 2 0. 4. Finally, concentrated sulphuric acid acts as an oxidizing agent. Sulphur and carbon, boiled in it, are oxidized: 2H 2 S0 4 + C -> 2S0 2 + 2H 2 + C0 2 and metals are converted into sulphates. The action on copper may be used for preparing sulphur dioxide. It should be noted that metals do not displace hydrogen, except from the dilute acid, SULPHUR TRIOXIDE AND SULPHURIC ACID 287 because the latter alone contains the hydrogen ions, which are thus discharged and liberated. The concentrated acid has usually to be heated, and then the molecules H 2 SO 4 give part of their oxygen for oxidizing purposes, and hydrogen is not liberated. Sulphates. A number of sulphates are in common use. Thus partly hydrated calcium sulphate 2CaSO 4 , H 2 is plaster of Paris, hydrated magnesium sulphate MgS0 4 , 7H 2 is Epsom salt, hydrated cupric sulphate CuS0 4 , 5H 2 O is blue-stone, used in bat- teries. The first is made by heating gypsum CaS0 4 , 2H 2 O, a common mineral. When the plaster of Paris is moistened with water, it "sets," and is thus used in making "plaster casts." The setting results from the reformation of gypsum: 2CaS0 4 , H 2 + 3H 2 -> 2CaSO 4 , 2H 2 O. Uses of Sulphuric Acid. The acid has innumerable applica- tions, some of which have already been mentioned. The dehy- drating power is valuable in making nitroglycerine and guncotton (pp. 382-384). As a cheap acid, it is used to clean iron before tinning and galvanizing. Under calcium phosphate its use in making fertilizers is explained (see p. 339). The unsaturated hydrocarbons, contained in natural petroleum, are oxidized by the air to dark-colored products. Kerosene and other oils, during their manufacture, are therefore agitated with concentrated sulphuric acid, which combines with such compounds, and leaves an oil which remains colorless when kept. Exercises. 1. Which contains more oxygen: (a) a phosphate or a phosphite; (b) a nitrite or a nitrate; (c) a borate or a per- borate? Name the acids corresponding to these six salts. 2. Make equations for: (a) the roasting of stannic sulphide (SnS2) giving SnO 2 ; (b) the action of concentrated sulphuric acid on silver giving silver sulphate and S0 2 (use partial equations); (c) the dissociation of sulphuric acid vapor. 288 TEXT BOOK OF CHEMISTRY 3. What are the formulae of magnesium sulphite and bisulphite, respectively ? 4. Give two reasons why boiling sulphuric acid, when spilt upon the flesh, causes most serious burns. 5. By what facts or tests could you recognize concentrated sulphuric acid ? 6. Why is the wooden laboratory shelf commonly " burned " where the sulphuric acid bottle stands? 7. What would be the reaction between sodium sulphite and sulphuric acid? 8. How could you distinguish between, and recognize, the sul- phide, sulphite, and sulphate of potassium? THE PERIODIC SYSTEM Classification, or the arrangement of facts on the basis of like- ness, is part of the method of science. It is needed to make possible the systematic description of the ascertained facts, which is a great aid to the memory, and to furnish a guide in investigation, by suggesting relations and so pointing out directions in which new facts of interest may be found. Thus, as an aid to memory, we have treated the halogens as one family. In each case, we have presented the properties common to all members of the group, and have then pointed out the differences. Metallic and Non-Metallic Elements. Thus far we have found the division into metallic and non-metallic elements very serviceable for classification in terms of chemical relations. This distinction we shall continue to employ. The metallic or positive elements (p. 158), (1) form positive radicals and ions containing no other element (cf. p. 174). Thus the metals give sulphates, nitrates, carbonates, and other salts, which furnish a metallic ion, such as Na + or K + , together with the ions S0 4 = , N0 3 ~, and COa". (2) Their hydroxides, KOH, Ca(OH) 2 , etc., give the same metallic ion, and the rest of the molecule forms hydroxide-ion. THE PERIODIC SYSTEM 289 That is to say, their hydroxides are bases and their oxides are basic. The metallic elements often enter, but only with other elements, into the composition of a negative ion, as is the case with manganese in K.Mn(>4, with chromium in K 2 .Cr 2 07, and with silver in K.Ag(CN) 2 . The non-metallic or negative elements (1) are found chiefly in negative radicals and ions. They form no nitrates, sulphates, car- bonates, etc., for they could not do so without themselves alone constituting the positive ion. We have no such salts of sulphur, carbon, or phosphorus, for example. (2) their hydroxides, al- though their formulae may be written C10 2 OH, P(OH) 3 , SO 2 (OH) 2 , furnish no hydroxyl ions, as this would involve the same conse- quence. These hydroxides are divided by dissociation, in fact, so that the non-metal forms part of a compound negative radical, and the other ion is hydrogen-ion, C1O 3 .H, PO 3 H.H 2 , S0 4 .H 2 . Their oxides are acidic. (3) Their halogen compounds, like PBr 3 and S 2 C1 2 (p. 267), are completely hydrolyzed by water, and the actions are not, in general, reversible. The halides of the typical metals are not hydrolyzed (see p. 356), and with those that are not typical, the action is reversible. The distinction is not perfectly sharp, however. Thus, zinc gives both salts like the sulphate, Zn.SO 4 , and chloride, Zn.Cl 2 , and compounds like sodium zincate NaHZnO 2 . Classification by Atomic Weights. Newlands (1863^) dis- covered a surprising regularity that became apparent when the ele- ments were placed in the order of ascending atomic weight. Omitting hydrogen (at. wt. 1) the first seven were: lithium (7), glucinum (9), boron (11), carbon (12), nitrogen (14), oxygen (16), fluorine (19). These are all of totally different classes, and include first a metal forming a strongly basic hydroxide, then a metallic element of the less active sort, then five non-metals of increasingly negative character, the last being the most active non-metal known. The next element after fluorine (19) was sodium (23), 290 TEXT BOOK OF CHEMISTRY which brings us back sharply to the elements that form strongly basic hydroxides. Omitting none, the next seven elements were sodium (23), magnesium (24.4), aluminium (27), silicon (28.4), phosphorus (31), sulphur (32), chlorine (35.5). In this series there are three metals of diminishing positiveness, followed by four non- metals of increasing negative activity, the last being a halogen very like fluorine. On account of the fact that each element resembles most closely the eighth element beyond or before it in the list, the relation was called the law of octaves. After chlorine the octaves become less easy to trace. That this periodicity in chemical nature is more than a coinci- dence is shown by the fact that the valence and even the physical properties, such as the specific gravity, show a similar fluctuation in each series. In the first two series the compounds with other elements are of the types : Lici, Gia, ECU, ecu, VT^rti T\/r~r^i Airn crr^i i PaOo, S02, Nad, MgCl 2 , Aids, Thus the valence towards chlorine or hydrogen ascends to four and then reverts to one in each octave. The highest valence, shown in oxygen compounds, ascends from lithium to nitrogen with values one to five, and then fails because compounds are lacking. In the second octave, however, it goes up continuously from one to seven. Again, the specific gravities of the elements in the second series, using the data for red phosphorus and liquid chlorine, are: Na 0.97, Mg 1.75, Al 2.67, Si 2.49, P 2.14, S 2.06, Cl 1.33. Mendelejeff's Scheme. In 1869 Mendelejeff published an important contribution towards adjusting the difficulty which the elements following chlorine presented, and developed the whole conception so completely that the resulting system of classification has been connected with his name ever since. Almost simultane- THE PERIODIC SYSTEM 291 ously Lothar Meyer made similar suggestions, but did not urge them with the same conviction or elaborate them so fully. The table on the following page, in which the atomic weights are ex- pressed in round numbers, is a modification of one of Mendelejeffs. The chief change from the arrangement in simple octaves is that the third series, beginning with potassium, is made to furnish material for two octaves, potassium to manganese and copper to bromine, and is called a long series. The valences fall in with this plan fairly well. Copper, while usually bivalent, forms also a series of compounds in which it is univalent. Iron, cobalt, and nickel cannot be accommodated in either octave, as their valences are always two or three. At the time Mendelejeff made the table, three places in the third long series had to be left blank, as a tri- valent element [Sc] was lacking in the first octave of the series, and a trivalent [Ga] and a quadrivalent one [Ge] in the second. These places have since been filled, as we shall presently see. The first two (the short) series have been split in the table, as lithium and sodium closely resemble potassium, while the remaining members of these series fall more naturally over the corresponding elements of the second octave of the third series. The fourth series (long) is nearly complete. It begins with an active alkali metal, rubidium, and ends with iodine, a halogen. The rule of valence is strictly preserved throughout the series, and in general the elements fall below those which they most closely resemble. The fifth, sixth, and seventh (long) series are incomplete, but the order of the atomic weights and the valence enable us satisfac- torily to place those elements which are known. The chemical relations to elements of the fourth series justify the position as- signed to each. Caesium, for example, is the most active of the alkali metals; barium has always been classed with strontium, and bismuth with antimony. In two cases a slight displacement of the order according to atomic weights is necessary. Cobalt is put before nickel because 292 TEXT BOOK OF CHEMISTRY 02 O2 OS OS Ss P* 52 as a? THE PERIODIC SYSTEM 293 it resembles iron more closely. Tellurium and iodine are placed in that order to bring them into the sulphur and halogen groups, respectively. Their valence and other chemical relations both require this. The general agreement, however, is very remarkable. General Relations in the System. In every octave the valence towards oxygen ascends from one to seven, while that towards hydrogen, in the cases of the last four elements (when they combine with hydrogen at all), descends from four to one. The physical properties fluctuate within the limits of each series in a similar way. The values of each physical constant for correspond- ing members of the successive series do not exactly coincide, how- ever. A progressive change, as we descend each vertical column, is the rule. Thus the specific gravities (water = 1) of the alkali metals rise from lithium (0.53) to caesium (1.87). In the same group the melting-points descend from lithium (186) to caesium (26.5). As yet no exact mathematical (quantitative) relation between the values for any property and the values of the atomic weights has been discovered; only a general (qualitative) relationship can be traced. Anticipating the discovery of some more exact mode of stating the relationship in each case, and remembering that similar values of each property recur periodically, usually at inter- vals corresponding to the length of an octave or series, the principle which is assumed to underlie the whole, the periodic law, is stated thus: All the properties of the elements are periodic functions of their atomic weights. That the chemical relations of the elements vary just as do the physical properties of the simple substances is easily shown. Thus, each series begins with an active metallic (positive) element, and ends with an active non-metallic (negative) element, the inter- vening elements showing a more or less continuous variation between these limits. Again, the elements at the top are the least metallic of their respective columns. As we descend, the members 294 TEXT BOOK OF CHEMISTRY of each group are more markedly metallic (in the first columns), or, what is the same thing, less markedly non-metallic (in the later columns; cf. p. 292). In the first series boron is the first non-metal we encounter. In the second series silicon is the first such element. In the third there is more difficulty in deciding. Titanium, vanadium, and germanium are usually, though with questionable propriety, classed as metallic elements.* Selenium is undoubtedly non-metallic. Arsenic is, on the whole, non-metallic. In the fourth series telu- rium is commonly considered to be the first non-metallic element. Thus a zigzag line, shown in the table, separates all the non-metallic elements from the rest of the elements, and confines them in the right-hand upper corner. A more compact form of the table is printed at the end of this book, opposite the rear cover. The only difference between this and the other is that the two octaves of each long series have been placed in the same set of seven main columns. The iron, palla- dium, and platinum groups occupy a column on the right of the main columns, and are often called collectively the eighth group. The newly discovered elements, found chiefly in the air, have been placed at the left-hand side. Since they do not enter into com- bination at all, their valence may appropriately be given as zero With the exception of argon, the values of their atomic weights agree well with this assignment. Hydrogen is the only common element whose place is still in debate. The valence is shown by the general formulae at the head of each column. Applications of the Periodic System. The system has found application chiefly in four ways: 1. In the prediction of new elements. Mendelejeff (1871) drew attention to the blank then existing between calcium (40) and tita- * In discussing chemical relations, the term metallic element is preferable to metal. The free element (e.g., arsenic) may have the luster of a metal and yet the element, in combination, may be non-metallic or negative. THE PERIODIC SYSTEM 295 nium (48). He predicted that an element to fit this place would have an atomic weight 44 and would be trivalent. From the nature of the surrounding elements, he very cleverly deduced many of the physical and chemical properties of the unknown element and of its compounds. In 1879 Nilson discovered scandium (44), and its behavior corresponded closely with that predicted. Mendele- jeff described accurately two other elements, likewise unknown at the time. In 1875 Lecoque de Boisbaudran found gallium, and in 1888 Winkler discovered germanium, and these blanks were filled. 2. By enabling us to decide on the correct values for the atomic weights of some elements, when the equivalent weights have been measured, but no volatile compound is known (cf. pp. 78 and 87). Thus, the equivalent weight of indium was 38 and, as the element was supposed to be bivalent, it received the atomic weight 76. It was quite out of place near arsenic (75), however, being decidedly a metallic element. As a trivalent element with the atomic weight 115, it fell between cadmium and tin. Later work fully justified the change. Quite recently, radium has been discovered, and found to have the equivalent weight 113 and to resemble barium. If, like barium, it is bivalent, it occupies a place under this element, in the last series. 3. By suggesting problems for investigation. The periodic sys- tem has been of constant service in the course of inorganic research, and has often furnished the original stimulus to such work as well. For example, the atomic weight of tellurium bore the value 128 when the table was first constructed, and it was confidently ex- pected that reexamination would bring this value below that of iodine (then 127, now 126.92). Several most careful studies of the subject have been made by different methods. It seems probable that the real value of the atomic weight is not far from Te = 127.5, and therefore more than half a unit greater than that of iodine. Since, however, mathematical correspondence is found nowhere in the system, the existence of marked inconsistencies like this need 296 TEXT BOOK OF CHEMISTRY not shake our confidence in its value when it is used with due con- sideration of the degree of correspondence to be expected. In the same way, incorrect values of many physical properties have been detected, and have been rectified by more careful work. 4. By furnishing a comprehensive classification of the elements, arranging them so as to exhibit the relationships among the physi- cal and chemical properties of the elements themselves and of their compounds. Constant use will be made of this property of the table in the succeeding chapters. Having disposed of the halogen and sulphur families (excepting the oxygen compounds of the former), situated, respectively, in the seventh and sixth columns of the table (at the end of this book), we shall presently take up nitrogen and phosphorus from the right side of the fifth column. Then from the fourth column, we shall select carbon and silicon, and from the third boron, leaving the other, more decidedly metal- lic elements for later treatment. Moseley's Atomic Numbers. We have seen that simple, mathematical relations between the atomic weights and the physi- cal or chemical properties of an element do not exist. In several instances, the atomic weights are not even in the same order as values of the properties. We have now obtained from another direction numbers which seem to be more fundamental even than atomic weights. Visible light, X-rays, and wireless electric waves are all vibra- tions of the same nature in the ether. They differ only in wave- length, the order of the wave-lengths being 10~ 5 cm., 10~ 8 cm., and 10 6 cm. (10 kilometers), respectively. Now, just as the spectrum of visible light is obtained by using a grating, on which the rulings are separated by distances of the order of the wave-length of such light, so ordinary crystals give spectra of X-rays, because they are composed of particles arranged in rows about one thousand times closer and so form a suitable grating for X-rays. This fact was first discovered by Dr. Luke of the University of Zurich (1912). THE PERIODIC SYSTEM 297 The X-rays are produced in an evacuated tube by cathode rays, which are streams of electrons emanating from the cathode (C, Fig. 88), when they strike the anticathode (A). With different elements on the anti-cathode, X-rays of slightly different wave-lengths, arid therefore giving different X-ray spectra, are produced. By using different elements, Moseley (1914) has found that the higher the atomic weight the shorter the wave-length of the character- istic X-rays. When the ele- ments are arranged in the order of these wave-lengths, whole numbers can be assigned to each which are inversely proportional to the wave-lengths of corresponding lines in their X-ray spectra. These atomic numbers have been determined for most of the elements, the atomic weights of which lie between those of aluminium and uranium. In the following table, the atomic numbers for these elements are given and, for the sake of greater completeness, numbers for the twelve elements preceding Al have been inserted also. ATOMIC NUMBERS (MOSELEY) FIG. 88. H 1 He 2 Li 3 Gl 4 B 5 C 6 N 7 O 8 F 9 Ne 10 Na 11 Mg 12 Al 13 Si 14 P 15 S 16 Cl 17 A 18 K 19 Ca 20 Sc 21 Ti 22 V 23 Cr 24 Mn 25 Fe 26 Co 27 Ni 28 Cu 29 Zn 30 Ga 31 Ge 32 As 33 Se 34 Br 35 Kr 36 Rb37 Sr 38 Y 39 Zr 40 Cb 41 Mo 42 - 43 Ru 44 Rh45 Pd 46 Ag 47 Cd 48 In 49 Sn 50 Sb 51 Te 52 I 53 Xe 54 Cs 55 Ba 56 La 57 Ce 58 Ta 73* W 74 - 75 Os 76 Ir 77 Pt 78 Au 79 Hg 80 Tl 81 Pb 82 Si 83 Po 84 - 85 Nt 86 - 87 Ra 88 Ad 89 Th 90 Ut 91 U 92 * The atomic numbers 59-72 are those of the metals of the rare earths: Pr 59, Nd 60, 61, Sa 62, Eu 63, Gd 64, Tb 65, Dy 66, Ho 67, Er 68, Tm 69, Yb 70, La 71, -72. It will be seen that there is a whole number available for every known element, up to and including uranium, and not omitting the rare elements which have no satisfactory place in the periodic 298 TEXT BOOK OF CHEMISTRY system. There are four blank numbers in the table, which corre- spond to three spaces below Mn in the periodic system, one under gold, and two more amongst the rare elements, indicating only six elements with atomic weights less than that of uranium yet to be discovered. The atomic numbers of argon and potassium place them in the chemically correct order, while the atomic weights do not. The same is true of cobalt and nickel and of tellurium and iodine. Finally, it is evident that the atomic weight of each ele- ment is, roughly, double its atomic number. The atomic numbers represent the number of unit positive charges of electricity in the nucleus of the atom of each ele- ment. Rutherford has shown that the nucleus contains almost the whole mass of the atom, although one or more electrons (negative) are present also. Thus, the positive nucleus of the hydrogen atom is 1700 times heavier than one electron. The nucleus, however, is very minute, having a diameter only about one-thousand million millionth of that of the whole atom. The atomic numbers apparently determine all the properties of each element, and are more fundamental than the atomic weights. The latter are secondary properties, in most cases modified by other factors, and in a few cases actually thrown out of order by such factors. Crystal Structure. In this connection it may be mentioned, that by using crystals of different substances as X-ray gratings, W. L. Bragg (1914) has been able to measure the distances be- tvveen the rows of particles in crystals. He also finds that the particles, the regular arrangement of which gives the structure (p. Ill) of the crystal (e.g., a cube of common salt), are not the molecules of the compound, much less aggregates of such mole- cules, but the atoms of the constituent elements. It would thus appear that the physical forces (if we may call them physical), which hold the crystalline solid together have subordinated the chemical, molecular structure out of existence, and have arranged THE PERIODIC SYSTEM 299 the constituent atoms, as the units of the structure, in a crystal- lographic pattern. Of course, when the crystal-form is destroyed, by melting, solution, or vaporization, the neighboring atoms re- main united in groups, constituting the chemical molecules of the substance. Exercises. 1. Can you explain the presence of free selenium in the flues of pyrite burners (p. 292) ? 2. How should you attempt to obtain H 2 Te, and what physical and chemical properties should you expect it to possess ? 3. Make a list of bivalent elements and criticize this method of grouping as a means of chemical classification. 4. Write down the symbols of the elements in the fourth series (that beginning with rubidium, and ending with iodine) on p. 292. Record the valence of each element toward oxygen, using for refer- ence the chapters in which the oxygen compounds are described. CHAPTER XXV OXIDIZING SUBSTANCES NITRIC acid, which we shall take up in the next chapter, shows a variety of very different chemical properties. One of these is the property of oxidizing other substances. In order to under- stand these oxidations, and to see at once the difference between the oxidizing actions, and the other actions which the acid exhibits, we must now learn more about oxidation in general. We can do this best through the study of three oxidizing substances which are of a simpler nature and are all in common use. These are ozone, hydrogen peroxide, and hypochlorous acid. OZONE O 3 . When oxygen is blown from a small nozzle through the tip of a Bunsen flame, a part of it is turned into ozone. The same thing happens when a platinum wire, heated by an electric current, is held under liquefied oxygen. This shows us that, to get ozone, we must add energy (for example, by strong heating) to oxygen. We learn, also, that the ozone must be cooled at once and kept cold. If it lingers in the cooler (but not cold) region round the flame, it decomposes again. Preparation of Ozone. In practice electrical energy, de- veloped by passing a "silent discharge" through the oxygen, is employed. This discharge gives ozone very easily, because its use involves no rise in temperature whatever. 30 2 ->20 3 . The poles of the induction coil are attached to the tinfoil upon the outside of the outer tube and the inside of the inner tube 300 OXIDIZING SUBSTANCES 301 (Fig. 80) . The " discharge " therefore passes through two layers of glass, as well as through the oxygen. The oxygen, from a cylinder of the gas, flows slowly through the space between the tubes. At FIG. 80. best, only 6 to ozone. into 7 per cent of the oxygen is usually changed i al Properties. Ozone (Greek, to sweW) is a gas of deep blue color, with a fresh, highly individual odor. It is more easily liquefied than is oxygen (b.-p. 119), and is also much more soluble in water. Its density is one-half greater than that of oxygen, and the formula 3 records this fact. Ozone and oxygen are, therefore, allotropic forms of the same ele- ment. Chemical Properties. Ozone, being produced by giving energy to oxygen, is at rather low temperatures (e.g., from 10 to 500) the less stable form of the element. Upon standing, and more quickly when warmed, it changes into oxygen, with liberation of the additional energy it contains. Being possessed of more internal energy than oxygen, ozone oxidizes the same substances as does oxygen, only more rapidly and vigorously. Hence, it instantly oxidizes sulphur dioxide to sulphur trioxide, and sulphurous acid to sulphuric acid: S0 3 + 2 . H 2 S0 3 + 3 302 TEXT BOOK OF CHEMISTRY For the same reason it oxidizes many substances not affected by ordinary oxygen. Thus, it rusts silver to black silver per- oxide Ag 2 2 . As a rule, only the one extra atom of oxygen in the molecule of ozone is used for the oxidation. Ozone also oxidizes a number of organic compounds which are unchanged in atmospheric oxygen. For example, when ozon- ized oxygen is bubbled through a dilute indigo solution, a yellow substance, of much paler tint, isatine, is formed, and the indigo is said to have been bleached. Indigo is taken, for illustration, because it is a most widely used dye, employed in dying navy-blue and blue-black goods, and is totally unaffected by light, and by oxygen, soap, and other ordinary substances: 20 3 -> 20 2 + 2C 8 H 5 NO 2 . indigo isatin Litmus, and the traces of coloring matter in wax, starch, flour, and ivory are all oxidized by it to colorless, or nearly colorless, substances. For this reason it is used commercially in bleaching the last-named materials. Ozone is sometimes recommended for use, in connection with ventilation, as a means of destroying minute organisms in the air. Recent investigations have shown, however, conclusively, that when thus diluted with air, it has little value as a germicide. It is employed by some cities (e.g., Paris, Florence, and St. Peters- burg) for sterilizing the water supply. From the foregoing we see that, by an oxidizing agent, like ozone, we ordinarily mean a substance which is a better oxidizer than free atmospheric oxygen. HYDROGEN PEROXIDE H 2 2 . Preparation. Sodium peroxide Na 2 O 2 , produced by burning sodium in dry air, can be dissolved, a little at a time, in ice-cold HYDROGEN PEROXIDE H 2 C>2 303 water. When this solution is acidified with hydrochloric or sul- phuric acid, a double decomposition takes place: 2HC1 -> 2NaCl + H 2 O 2 and a dilute solution of hydrogen peroxide (mixed with common salt) is obtained. The nature of the action shows the product to be an acid, with the negative radical 62". For manufacturing purposes it is more convenient to use barium peroxide BaO 2 , suspended in water, and sulphuric acid: BaO 2 + H 2 SO 4 -> BaSO 4 I + H 2 O 2 because the barium sulphate is insoluble and can be removed by filtration. In pharmacy a 3 per cent solution is the one commonly sold. As hydrogen peroxide decomposes at 100, the pure substance can be obtained by reducing the pressure of the air in the closed distilla- tion apparatus to 26 mm., and distilling off the water at 27. Properties. Hydrogen peroxide is a colorless, syrupy liquid of sp. gr. 1.5, misdble with water in all proportions. Dilute solutions have a metallic taste. Chemical Properties. 1. As an add, hydrogen peroxide enters into double decomposition, particularly with bases, giving salts containing the radical O*: Ca(OH) 2 + H 2 O 2 -> Ca0 2 i + 2H 2 O. The salts are now often called peroxidates, instead of peroxides, as sodium peroxidate and barium peroxidate, and the acid itself is sometimes called hydrogen peroxidate. The latter is a very feeble acid, hardly affecting litmus. 2. When the solution is heated, the compound decomposes, with evolution of heat, giving water and oxygen : 2H 2 2 -+2H 2 + O 2 t- 304 TEXT BOOK OF CHEMISTRY Contact agents hasten the decomposition. Thus, it takes place with frothing when the cold solution is applied, as an antiseptic, to cuts or sores, or when powders, such as manganese dioxide, are thrown into the solution. 3. Since hydrogen peroxide, like ozone, gives off oxygen with liberation of energy, it is an oxidizing agent also. In this respect its behavior is very similar to that of ozone. Thus, it oxidizes sulphurous acid to sulphuric acid: H 2 S0 3 + H 2 2 -> H 2 S0 4 + H 2 0. It oxidizes colored organic compounds to colorless ones, and is, therefore, used in bleaching hair, feathers, silk, and ivory. It is also fatal to micro-organisms, and is, therefore, employed in medicine to disinfect wounds and as a throat wash. In oil paintings the high lights are produced in part with white lead (carbonate of lead) . These disappear and the picture darkens with age, because hydrogen sulphide in the air turns the white lead into the black lead sulphide, PbS. In picture restoring the latter is oxidized to lead sulphate, which is white, by treatment with hydrogen peroxide solution. PbS + 4H 2 2 -> PbS0 4 + 4H 2 0. 4. A beautiful reaction is used as a test. When potassium dichromate solution is acidified with sulphuric acid, to liberate the dichromic acid: K 2 Cr 2 7 + H 2 S0 4 - H 2 Cr 2 7 + K 2 S0 4 and a drop of the mixture containing the latter acid is added to aqueous hydrogen peroxide, a substance possessing a deep, brilliant blue color is formed. By this test the presence of hydrogen per- oxide in rain water can often be demonstrated. HYPOCHLOROUS ACID HC1O. Pure Hypochlorous Acid. A pure solution of the acid may be made by dissolving the corresponding anhydride in water. The anhydride, chlorine monoxide C1 2 O, is a brownish-yellow, HYPOCHLOROUS ACID HClO 305 explosive gas, made by passing chlorine gas over warmed mercuric oxide * (Fig. 46, p. 134) : 2C1 2 + HgO -> HgCl 2 + C1 2 0. C1 2 + H 2 - 2HC10. As an acid, hypochlorous acid is very weak, being very little decomposed into its ions, H + and (C10)~. It is unstable, exposure to sunlight being sufficient to cause it to give up oxygen, which rises in bubbles through the solution (Fig. 81) : 2HC10->2HC1 + 2 t Heat is given out in the action, and the stable hydrochloric acid remains. It is a most active oxidizing agent, because of this tendency to give up oxygen with libera- tion of energy. Thus, its solution instantly oxidizes sulphurous acid to sulphuric acid: H 2 S0 3 + HClO -> H 2 S0 4 + HC1. FlG> 81 ' The solution also oxidizes organic colbred substances, producing colorless or less strongly colored ones: C 16 H 10 N 2 2 + 2HC10 -> 2C 8 H 5 N0 2 + 2HC1. indigo isutin Used as a disinfectant (see p. 308), it oxidizes and destroys bac- teria. Hypochlorous acid is more energetic as an oxidizing agent than is ozone or hydrogen peroxide. Chlorine-Water. Hypochlorous acid made from chlorine monoxide is much too expensive for application in the bleaching industry. But fortunately, on account of its great activity as an oxidizer, a very dilute solution is all that is required, and is, indeed, * The mercuric oxide must be made by precipitation from sodium hydroxide and mercuric nitrate solutions, and must be well washed with water, and completely dried before use. 306 TEXT BOOK OF CHEMISTRY all that can safely be used. Hence, in practice, chlorine-water, which can be made very cheaply, is the form in which hypo- chlorous acid is always employed. It will be recalled that chlorine acts chemically upon water (p. 139): C1 2 + H 2 <= HC1 + HC1O giving hydrochloric acid and hypochlorous acid. The action is reversible (read the equation backwards), and in half -saturated chlorine solution about one-third only of the chlorine has under- gone the change shown in the equation. But, if a substance which can be oxidized, such as a dye (attached, perhaps, to cloth), is introduced into the solution, the HC1O which is present transfers its oxygen to the dye-stuff. This leaves HC1 alone in the solution, and stops the backward reaction. Hence more of the chlorine acts upon the water, and more hypochlorous acid is formed. This, in turn, is used up. Thus, in a few moments, all the free chlorine is gone, only dilute hydrochloric acid remains, and a colorless organic compound is left on the cloth or in the solution. The mechanism seems complicated, but in reality is of a sort with which we are quite familiar. It is like keeping one's money in a bank, where it exists, perhaps, largely in the form of loans and bonds. We then draw it out as it is needed, instead of carrying it all, loose, in our pockets. Sunlight acts upon chlorine-water, exactly as it does on any other solution of hypochlorous acid, causing oxygen to be given off. In this case, also, only hydrochloric acid finally remains: 2C1 2 + 2H 2 < 2HC1 + 2HC10 - 2HC1 + 2 | Chlorine itself is often, erroneously, spoken of as the bleaching agent. If a dry, colored cloth be hung for a week in chlorine, dryed by having sulphuric acid in the bottle (Fig. 82), little or no change FIG. 82. HYPOCHLOROUS ACID HClO 307 in color will occur. But a wet rag is bleached as soon as the chlo- rine has time to dissolve in the water, and give the necessary hypo- chlorous acid. Bleaching. Cotton and linen, in their original states, are not pure white. Bleaching is therefore an extensive and most impor- tant industry. The yarn or cloth must first be freed from cotton- wax and tannin, since the former would hinder the action of the bleaching agent, and both would also make the subsequent dyeing uneven. The material is therefore first boiled with dilute caustic soda-solution, and washed with water. The cloth is then satu- rated with bleaching powder solution, and is piled loosely until the coloring matter has been oxidized. The goods are then washed with extreme thoroughness. Bleaching powder CaCl(OCl) is made by the action of chlo- rine on quicklime: Cl.->CaCl(Oa). A strong acid, such as sulphuric acid, acts upon it by double decomposition: CaCl(OCl) + H 2 S0 4 *4 CaS0 4 + HC1 + HClO <= C1 2 + H 2 0. A weak acid, however, like carbonic acid (see p. 201), can liberate the hypochlorous acid only. In bleaching, as a rule, no active acid is added. The effect is produced by the bleaching powder itself, supplemented by hypo- chlorous acid liberated by the action of carbon dioxide from the air. The final washing, to remove all traces of chlorine and bleach- ing powder, is absolutely necessary. If not removed, the hypo- chlorous acid acts gradually upon the cotton or linen, and "rots" it. Bleaching agents, when used in the household, carelessly, are liable to cause extensive damage from this cause. A dilute solu- tion of sodium thiosulphate (photographers' "hypo") is often 308 TEXT BOOK OF CHEMISTRY used, as "antichlor," to interact with and remove the last traces of chlorine. Cotton and linen (C 6 Hio0 5 )y are rather indifferent chemical substances (p. 227), and stand brief contact with dilute chlorine- water without much alteration. But wool and silk contain com- pounds of nitrogen (proteins) and are acted upon as rapidly as are the traces of colored matter themselves. Hence sulphurous acid is used for bleaching these materials. Bleaching Powder as a Disinfectant. A disinfectant is a substance which destroys bacteria and other minute, and often harmful organisms. Bleaching powder CaCl(OCl) has a distinct odor of hypochlorous acid (or C1 2 0, not chlorine). This is due to the action of carbonic acid, derived from carbon dioxide in the air, which is able to displace the weak acid (HC10), but not the much stronger one (HC1) : 2CaCl(OCl) + H 2 C0 3 -CaCl 2 + CaC0 3 + 2HC10. The dry powder will disinfect the surrounding air, because of the oxidizing power of the hypochlorous acid thus liberated. When an epidemic of typhoid fever occurs, it is usually traced to the presence of colon bacilli and typhoid organisms in the drink- ing water. The most effective means of destroying these bacilli is to add, at the distributing point, a small proportion of bleaching powder (about 20 pounds per million gallons of water). The salt is hydrolyzed (p. 356), so far as the hypochlorous acid radical is concerned, giving basic calcium chloride, and the free acid kills the bacteria: CaCl(OCl) + H 2 O^CaCl(OH) + HC10. Oxidations Previously Mentioned. The simplest oxida- tions are the unions of oxygen with metals and with non-metals : 2 ->2CO. O 2 -S0 2 . HYPOCHLOROUS ACID HClO 309 The union of more oxygen with an oxide or an oxacid is oxidation also: 2 ->2CO 2 . 2 -+2H 2 S0 4 . The removal of hydrogen from hydrogen chloride, for the purpose of preparing chlorine, was (p. 137) likewise defined as oxidation: 2 + 4HC1 - 2H 2 O + 2C1 2 . 2KMn0 4 + 16HC1 -> 8H 2 O + 2KC1 + 2MnCl 2 + 5C1 2 . Every oxidation is accompanied by a reduction of the oxidizing agent. Thus, in the above instances, the free oxygen is reduced to water. The permanganate is also reduced, as we can see by the fact that its oxygen has been removed. The appearance of a product that could only be formed by reduction is sometimes the first thing that calls our attention to the fact that an oxidizing action has occurred. When concentrated sulphuric acid acts upon carbon (p. 286), the C0 2 given off shows that there was oxidation, but the odor of the sulphur dioxide is the first thing we notice when doing the experiment: 2H 2 S0 4 + C -> C0 2 + 2H 2 + 2S0 2 . Removal of the elements of water, for example, from an oxacid like carbonic acid, is neither oxidation nor reduction, for hydrogen and oxygen are both removed: H 2 SO 3 -SO 2 + NH 4 OH->NH 3 The examples in the present chapter introduce nothing new. Ozone-, hydrogen peroxide, and hypochlorous acid are oxidizing agents, like potassium permanganate. But the actions are simpler than are those of the latter substance. We can now see that oxidation, in these cases, consists always in adding oxygen or removing hydrogen. 310 TEXT BOOK OF CHEMISTRY Other Cases of Oxidation. But oxygen is only one of a class of elements which we call non-metallic or negative elements, so that we do not restrict the term "oxidation" to actions involving oxygen. Thus, forming a chloride, by throwing a metal into chlorine, is oxidation also: 3Cl 2 -2SbCl 3 . Similarly, changing ferrous chloride FeCl 2 to ferric chloride Feds is oxidation: Cl 2 -2FeCl 3 . In every compound one of the elements is relatively positive and the other relatively negative. Iron is positive, and chlorine negative. In carbon monoxide carbon is (relatively) positive and oxygen negative. In ferrous sulphide iron is positive and sulphur (relatively) negative. Oxidation, then, is introducing, or increasing the proportion of the negative element, or removing, or reducing the proportion of the positive element (see also pp. 134, 135). Reduction is the converse. Oxidation and Valence. Combining oxygen or chlorine with a metal raises the active valence of the latter from zero to some finite value. Metallic copper has no valence in use. In CuCl2 the copper is employing the valence II. The copper has been oxidized. Similarly, changing FeCl 2 to FeCl 3 increases the active valence of the iron from II to III. Conversely, changing HC1 to C1 2 alters the active valence of the chlorine from I to zero. Hence, oxidation may be defined as increasing the active valence of a positive element or decreasing that of a negative element. Reduction is the converse. Oxidation and Electrons. Finally, since increasing the valence of a negative atom means adding one or more electrons to that atom, and increasing the positive valence of an atom means HYPOCHLOROUS ACID HClO 311 removing one or more electrons, we reach the briefest definition by saying: Oxidation is removing electrons and reduction is adding electrons. Energy and Chemical Action. It will have been noticed that there are two kinds of chemical actions: (1) those which, once started, proceed of themselves, and (2) those which have to be forced along by application of heat or some other form of energy. To the former kind belong most of the actions we use, such as combustion, displacements, and ordinary oxidation. These are accompanied by the liberation of energy in some form, usually heat. To the second class belong the decomposition of mercuric oxide, the electrolysis of salts and the formation of ozone. They are accompanied by absorption of energy, and they come to a stop when we cease supplying the energy. Further than this, the actions which proceed of themselves most rapidly are generally those which give out most energy. Finally, if free oxygen will not oxidize silver or indigo, we can accomplish the result by using a substance, like ozone or hypo- chlorous acid, capable of giving up energy when it liberates oxygen. The two lots of energy, being set free together, make a total libera- tion of energy sufficient to carry the whole action through success- fully. The chemically most active metals and oxidizing agents, there- fore, are those which, in general, contain most available energy, and are therefore capable of liberating the most energy when they enter into chemical action. Exercises. 1. Why does air containing ozone lose the latter (by change into oxygen) quicker when warm than when cold? 2. What other element, besides oxygen, have we found to occur in allotropic forms? Was the energy content of these forms different? 3. Make equations for: (a) the oxidation of carbon monoxide 312 TEXT BOOK OF CHEMISTRY to carbon dioxide by ozone; (b) the action of the components of chlorine-water upon calcium hydroxide to give bleaching powder. 4. Mark the valences of the radicals in barium peroxide and hypochlorous acid. 5. By what experiments can we ascertain whether an acid is strong or weak? 6. What volume (at and 760 mm.) of oxygen would be obtained by the decomposition of the hydrogen peroxide in 1 kilo- gram of the 3 per cent solution? 7. Are the following to be classed as oxidations: (a) the pro- duction of glucose from starch; (6) the action of heated cupric oxide on ammonia; (c) the action of magnesium on ammonia? If not, what is each? CHAPTER XXVI NITRIC ACID NITRIC acid HN0 3 is used in large quantities for making- explo- sives like guncotton, plastics like celluloid (Chap. XXXIII), as well as innumerable drugs and dyes. Nitrates are largely used as fertilizers (p. 320). Manufacture. Nitric acid is obtained in two ways, namely, by the action of sulphuric acid upon natural sodium nitrate and, to a smaller extent, by oxidation of the nitrogen of the atmosphere. The latter process will be referred to in a later section. Sodium nitrate, Chile saltpeter, is found in an immense deposit (2 by 220 miles) on the boundary of Chile and Peru. This salt is mixed with concentrated sulphuric acid hi iron retorts and gently heated to drive off the nitric acid. The sodium-hydrogen sulphate remains in the retort: NaNO 3 + H 2 S0 4 <=* NaHS0 4 + HN0 3 T The vapor is condensed in glass tubes (cooled with water) and the acid collected in vessels of earthen- ware. Sulphuric acid (b.-p. 338) is used because it is much less volatile than nitric acid, and so only the latter is vaporized. The acid boils at 86 FIQ ^ (760 mm.), but, to prevent loss by decomposition, a lower boiling-point is secured by reducing the pressure in the whole apparatus. In the laboratory the same action is employed, without, however, the reduction in the pressure (Fig. 83). 313 314 TEXT BOOK OF CHEMISTRY Physical Properties. Pure nitric acid is a colorless liquid, boiling at 86. It is miscible in all proportions with water. The vapor, like hydrogen chloride gas, condenses moisture from the air, giving a fog of droplets of the solution. The concentrated nitric acid of commerce contains 68 per cent of the acid (32 per cent Aq) and boils at 120.5. Points About Oxacids. In dealing with an oxacid, there are some things which we must acquire the habit of keeping in mind. Thus, an oxacid, as we have seen, can be deprived of water, leaving the anhydride. The chemist always thinks of the one as soon as the other is mentioned. If the acid is named, he instantly subtracts water from its formula to get the formula of the anhy- dride : H 2 C0 3 -H 2 + C0 2 . Skeleton: HC10 -> H 2 O + ? To balance the second equation, hydrogen atoms must be supplied in pairs, so the HC1O must be multiplied by two: Balanced: 2HC10 U H 2 + C1 2 O. The beginner must beware of making an arithmetical error either in multiplying HC10 by 2 (which gives H 2 C1 2 2 ), or in then sub- tracting H 2 0. Absurd as it may seem, he is likely not to get C1 2 O, but some incorrect formula. With nitric acid the operation is similar : Skeleton: HNO 3 -> H 2 + ? Balanced: 2HN0 3 -4 H 2 O + N 2 5 . This operation should be practiced: Skeleton: HC10 4 -> H 2 O + ? H 3 PO 4 -> H 2 + "? Balanced: 2HC10 4 - H 2 O + ? 2H 3 PO 4 -* 3H 2 O + ? The valence of the characteristic non-metal is ascertained by this process. What is the valence of Cl in HC10? The anhydride is NITRIC ACID 315 C1 2 O. The valence is Cl 1 . What is the valence of S in H 2 SO 4 ? The anhydride is S0 3 . The valence is S VI . What is the valence of N in HN0 3 ? The anhydride N 2 O 5 shows the valence to be N v . What are the valences of Cl in HC10 4 and of P in H 3 P0 4 and in HPO 3 ? When we get S0 2 from sulphuric acid, by a chemical action, how do we know that the acid has been reduced (and something else oxidized) ? Because in sulphuric acid we have S VI 03 n , and in the product S IV 2 n . The valence of S has been lowered from VI to IV. When from nitric acid we get N 2 C>5, has there been reduction? No, because N 2 5 -f H 2 O = H 2 N 2 O 6 = 2HNO 3 (nitric acid). If we get N0 2 or NO, has there been reduction of the acid? Yes, because the valence has been reduced: N^Ojj, N n O. We then proceed to pick out the other substance that has been oxidized. Analyzing the formula of the acid, to get that of the anhydride, also aids us to balance equations, as we shall presently see. Chemical Properties Decomposition of Nitric Acid: Equation Making. The acid is not very stable. It decom- poses, in part, even when simply distilled (86), giving a red gas, nitrogen peroxide N0 2 , and oxygen: Skeleton: HNO 3 -> N0 2 + 2 + ? We perceive at once that water must be formed as well: Skeleton: HNO 3 -+ N0 2 + 2 + H 2 O. Balancing (partially) : 2HNO 3 -> 2NO 2 + O 2 + H 2 O. This gives us H 2 O, correctly, but we have still only 6O on the left and 7O on the right: Balanced: 4HN0 3 -+ 4N0 2 + 2 + 2H 2 O. The operation becomes much simpler if we write the formula of the acid so as to show the anhydride: 2HNO 3 ( = H 2 O,N 2 O 5 ) -> 2NO 2 + O + H 2 O. 316 TEXT BOOK OF CHEMISTRY We now see that each N 2 5 will give 2N0 2 + 0, and to get 0, we must double the whole: 4HN0 3 ->4N0 2 + 2 + 2H 2 0. This is the decomposition when the acid is heated by itself. When a reducing agent is present, oxides containing less oxygen than N0 2 may be formed. The oxides usually obtained are the oxide N0 2 and nitric oxide NO. Nitric Acid as an Acid. Nitric acid turns blue litmus red, and is a very active, highly ionized acid. With bases it gives nitrates, which can be obtained from the solution by evaporation. NaOH+ HN0 3 -H 2 + NaN0 3 . Ca(OH) 2 + 2HN0 3 -+2H 2 + Ca(NO 3 ) 2 . Nitric Acid as an Oxidizing Agent. Since nitric acid gives up oxygen with liberation of energy, it is also active as an oxidizing agent. Glowing charcoal, as a powder or in the form of a FIG. 84. stick, will burn when pure nitric acid is poured upon it (Fig. 84). Carbon dioxide and red nitrogen per- oxide NO 2 are evolved. Skeleton: HN0 3 + C -> N0 2 + C0 2 + H 2 0. Analyzed: 2HN0 3 (H 2 0,N 2 5 ) + C -> N0 2 + CO 2 + H 2 . Each N 2 5 (from 2HNO 3 ) will give 2N0 2 + O. For CO 2 , 20 is required, and therefore 4HN0 3 is needed: Balanced: 4HN0 3 + C -> 4N0 2 | + CO 2 | + 2H 2 0. Nitric acid oxidizes indigo and other colored organic compounds, in the same way as do the three oxidizing agents described in the preceding chapter. It also oxidizes hydrochloric acid, upon which hydrogen peroxide does not act, so that it is a more active oxidizing agent than is that substance : HNO 3 + 3HC1 -> NOC1 + 2H 2 O + C1 2 . NITRIC ACID 317 The mixture of concentrated hydrochloric acid, nitric acid, and water is called aqua regia, and has strong oxidizing properties, due to the presence of hypochlorous acid (from C1 2 -h H 2 0) as well as nitric acid. The NOC1 + H 2 O gives HN0 2 + HC1. Action of Nitric Acid on Metals. Magnesium, and metals above it in the activity list, will displace hydrogen freely, especially from diluted nitric acid: Mg + 2HN0 3 -> Mg(N0 3 ) 2 + H 2 | But, with the less active metals, oxidation takes place, and instead of hydrogen, we get water and, of course, a reduction product of the nitric acid. The nitrate of the metal, however, is formed also. Even metals, like copper and silver, which do not displace hydro- gen, are acted upon by nitric acid in the same way. For example, diluted nitric acid acts vigorously upon copper, giving nitric oxide NO as the reduction product: Skeleton: Cu + HN0 3 -> Cu(NO 3 ) 2 + NO + H 2 0. If hydrogen could be displaced, the equation would be: Partial (1): Cu + 2HN0 3 -> Cu(N0 3 ) 2 + (2H) . But nitric acid turns this to water instead, and each N 2 0s gives 2NO -{- 3O, or enough oxygen to oxidize 6H: Partial (2) : 2HN0 3 ( = H 2 0,N 2 O 5 ) + 6H -> 3H 2 + 2NO. As hydrogen is not a product, and must cancel out, we now multi- ply Partial (1) throughout by 3, and add (1) and (2) : Partial (1) X 3: 3Cu + 6HNO 3 -* 3Cu(N0 3 ) 2 + (6H). Partial (2): (6H) + 2HNO 3 -+4H 2 + 2NO. Balanced: 3Cu + 8HNO 3 -> 3Cu(NO 3 ) 2 + 4H 2 + 2NO. The nitric oxide is a colorless gas, but unites with oxygen in the air to give NO 2 . 318 TEXT BOOK OF CHEMISTRY A test for a nitrate may be founded on this action. To the nitrate sulphuric acid is added to liberate nitric acid. Then copper turnings are thrown in to give NO. A gas turning red as it meets the air shows that a nitrate was present. Action Upon Organic Compounds. Nitric acid stains the skin and nails yellow, by giving colored compounds with the pro- teins. It gives similar compounds with wool (a protein), and therefore produces, on clothing, yellow stains which cannot be removed. The explosives made by the use of nitric acid are discussed in Chap. XXXIII. Nitric Oxide NO. This oxide is made by the action of diluted nitric acid upon copper, by the action already discussed (p. 317). It is a colorless gas, and almost insolubk in water. Vigorously burning phosphorus continues to burn in it: 4P + 10NO - 2P 2 5 + 5N 2 . Its most important property is that of uniting with oxygen to give nitrogen peroxide : 2NO + 2 -2N0 2 . Nitrogen Peroxide NO 2 . This oxide is formed by the union of oxygen with nitric oxide. It is given off, also, when concen- trated nitric acid acts upon metals and other reducing substances. It is further produced, along with oxygen, when nitrates, excepting those of potassium and sodium are heated, dry: 2Cu(N0 3 ) 2 -> 2CuO + 4N0 2 + 2 . Potassium and sodium nitrates, when heated, give off only oxygen, and leave the nitrites: 2KN0 3 ->0 2 + 2KN0 2 . NITRIC ACID 319 Nitrogen peroxide is a red gas. When cooled, it becomes pale yellow, and its density becomes twice as great. This is due to the formation of molecules of the formula N 2 04: The most interesting property of nitrogen peroxide is its action upon water, whereby nitric acid is formed, and nitric oxide escapes : 3N0 2 + H 2 - 2HNO 3 + NO. When oxygen is present also, then the NO gives more NO 2 , and this in turn gives more nitric acid. This action plays an important part in the making of nitric acid from the nitrogen of the air (see next section). The gas is sometimes used for bleaching flour, but traces of the oxide remain in the bread. Fixation of Atmospheric Nitrogen. Oxygen and nitrogen have no natural tendency to combine at the ordinary temperature, but rather the reverse the compounds tend to decompose. But a high temperature will supply the necessary energy. Even so, however, the union extends to only 1 per cent of the mixture at 2000 and 5 per cent at 3000: N 2 + O 2 <=* 2NO. Nevertheless, owing to the cost, and limited supply of natural nitrates, machinery has been devised, and is now in successful use, for carrying on the combination on a commercial scale. Three devices are in use, and all employ hydro-electric power. In the Birkeland-Eyde process (Fig. 85), used at Notodden and elsewhere in Norway, an arc discharge between rods of carbon is spread, by the influence of powerful electromagnets, into a circular FIG. 85. 320 TEXT BOOK OF CHEMISTRY WATCH foa\ brush discharge several feet in diameter. The figure shows a cross section of the space filled by the discharge. In the center is a section of one of the carbon rods. Air is blown through the flame, giving 1 per cent of NO, and is cooled to permit of union of the nitric oxide with oxygen, to give the peroxide, N0 2 . The air containing NO 2 is then passed through absorbing towers down which water trickles. Here the action mentioned in the last section takes place, and an aqueous solu- tion of nitric acid is produced. The yield is 70 g. of nitric acid per kilowatt-hour, and the net earnings in 1911 were $350,000. The nitric acid is mixed with calcium hydroxide (slaked lime) : Ca(OH) 2 + 2HN0 3 - Ca(NO 3 ) 2 + 2H 2 to give calcium nitrate, which, being very solu- ble, is sold for use as a fertilizer. The Badische process, used in the same factories in Norway, employs a discharge through a tube 22 feet long (Fig. 86). The column of air rotates as it traverses the tube and so every part is exposed to the discharge. The Pauling process, used at Gelsenkirchen, Germany, and Nitrolee, So. Carolina, uses pre- heated air, and a different arrangement of the discharge. The principles employed are, how- ever, the same. Other reactions involving the fixation of atmospheric nitrogen are discussed under calcium cyanamide (p. 435) and "root nodules" (p. 452). Nitrous Acid HNO 2 and Nitrous Anhydride N 2 O 3 . When an acid, such as sulphuric acid, is added to a solution of \Exir \ENTRMCe FIG. 86. NITRIC ACID 321 a nitrite, like potassium nitrite (p. 318), nitrous acid HN0 2 is formed : 2KNO 2 + H 2 SO 4 -> K 2 S0 4 + 2HN0 2 . Nitrous acid, however, like sulphurous acid, is unstable and nitrous anhydride N 2 O 3 is at once liberated and escapes as a gas: 2HN0 2 -H 2 This gas is used as a catalytic agent in the chamber process (p. 281) for making sulphuric acid. The acid and its anhydride are employed in the manufacture of many dyes. Nitrous Oxide N 2 O. When ammonium nitrate, a white salt, is heated, it decomposes into steam and nitrous oxide: 2H 2 |+ N 2 | . Nitrous oxide is somewhat more easily liquefied (b.-p. 90) than is carbon dioxide. At 12 its vapor pressure is 41 atmospheres. It is sold, as a liquid, in steel cylinders, and used as an anaesthetic for minor operations, chiefly in dentistry. The hysterical symp- toms which accompany its use caused it to be named "laughing gas." Like oxygen, it relights a glowing splinter of wood, and supports combustion brilliantly. It does not interact with nitric oxide (p. 318), as does oxygen, however, to form nitrogen peroxide. Exercises. 1. What is the valence of carbon in carbonic acid H 2 CO 3 ? 2. What is the anhydride of nitrous acid HNO 2 , and what the valence of nitrogen in this compound? 3. When nitric acid acts upon copper, which substance is oxi- dized and which reduced? State the effect upon the copper in terms of electrons. 4. How could you show experimentally that both nitrogen per- oxide and oxygen are formed when cupric nitrate is heated ? 322 TEXT BOOK OF CHEMISTRY 5. How could you distinguish nitric oxide, (a) from hydrogen, (6) from oxygen? 6. How could you distinguish nitrous oxide from oxygen? 7. Make a list of the names and formulae of the oxides of nitro- gen, arranging them in the order of increasing proportions of oxygen. 8. How should you make ammonium nitrate (p. 316)? CHAPTER XXVII THE HALOGEN FAMILY THE elements, if we may judge from those studied or mentioned thus far, may be divided into two classes the metallic or positive elements, like sodium, zinc, iron and magnesium, and the non- metallic or negative elements, like oxygen, sulphur, nitrogen and chlorine. The former give positive ions Na + , Zn++, Fe" 1 " 1 ", Mg ++ . The latter do not behave as positive ions. In other words, although we have nitrates and sulphates of all the metals, we have no nitrate or sulphate of oxygen or nitrogen. Natural Families of Elements. We have a simple means of subdividing within each of these two classes. We can place together the elements of like valence. Thus Na + and K + or Zn++ and Mg++ resemble one another in two respects metallic char- acter and valence. Also, = and S = form one group and Cl~, Br~, I~ and F~ form another. Groups of this kind are often spoken of as natural families of elements. The last group is called the halo- gen family, from the Greek for salt-producing, because these ele- ments combine with sodium to give substances all resembling common salt. Usually, the elements of one family resemble one another in a number of ways, and show at the same time a grada- tion in properties which it is interesting to study. The Halogen Family. The halogens are all negative ele- ments, and univalent. They show a gradation in properties which follows the order of their atomic weights. Here are a few properties, as examples: 323 324 TEXT BOOK OF CHEMISTRY Element. At. Wt. State. B.-P. Color. Solubility (15) in 1 c.c. Aq. Heat of Formation K Com- pound. Fluorine 19.0 gas -187 yel. (liq.) 118,100 Chlorine 35 5 gas 34 or.-yel. 072 104,300 Bromine 79 9 liq. 59 (Hq.) brown 0.032 95,100 Iodine 126 9 solid 184 violet 0.00015 80,100 It will be seen that, as the atomic weight increases, the boiling- point (b.-p.) rises, the color deepens, the solubility diminishes, and the heat of union with one atomic weight potassium becomes smaller. The last, and the atomic weight, are chemical properties, and they place the elements in the same order as do the physical properties.* BROMINE Br. The element was discovered by Balard in 1826 and derives its name from its offensive odor (Greek, a stench). Preparation. The salt deposits and natural salt wells of Cheshire, of Germany, and of Michigan, West Virginia, Ohio, and Connecticut, contain some bromides, along with large quantities of common salt. When the latter has been largely separated by evaporation and crystallization, the bromides of sodium and magnesium, which are more soluble, collect in the mother liquor. The bromine can be liberated at the positive electrode by elec- trolysis. But usually a chemical process is employed. In Germany chlorine gas is dissolved in the liquor. This displaces the bromine, and the latter can be distilled out by heating: 2NaBr + C1 2 -> 2NaCl + Br 2 . 2Br~ + C1 2 -> 2C1- + Br 2 . * For a complete table of all the families of elements, see the table of the Periodic System, facing the table of atomic weights (inside rear cover). BROMINE Br. 325 In the United States oxidation by pulverized manganese dioxide and sulphuric acid is generally employed, and this method can be used in the laboratory (Fig. 83, p. 313). The two reagents give manganous sulphate and oxygen: Partial (1): H 2 S0 4 + Mn0 2 -> MnS0 4 + H 2 -f (0). Partial (2) X 2 : 2H 2 S0 4 + 2NaBr - 2NaHS0 4 + 2HBr. Partial (3): (0) + 2HBr -+ H 2 O + Br 2 . Balanced: 3H 2 S0 4 + Mn0 2 -f 2NaBr - MnSO 4 + 2NaHSO 4 + 2H 2 + Br 2 . Physical Properties. Bromine is a liquid of a deep red-brown color and the vapor, of the same color, has a suffocating odor. It boils at 59. It is moderately solubk in water, giving a 3.2 per cent solution (bromine-water), and is very soluble in carbon di- sulphide. The density of the vapor gives it the formula Br 2 . Great care must be used in handling bromine, as, when spilt upon the skin, it kills the tissues and the sore is very liable to become infected. Treatment of Burns. Burns made by bromine or strong acids should be washed instantly with water and then with bi- carbonate of soda solution, and covered thickly with vaseline, or a salve of boric acid in lanoline, to protect them from infection. Chemical Properties. A jet of burning hydrogen will con- tinue to burn in bromine vapor, giving hydrogen bromide, which fumes in moist air like hydrogen chloride : H 2 + Br 2 ->2HBr. Many of the metals, when thrown in the form of powder, leaf, or foil, into bromine vapor, combine directly, giving bromides. The action is similar to that with chlorine, but less vigorous. Hydrogen Bromide HBr, Preparation. Hydrogen and bromine vapor unite rather slowly, and this direct union could at 326 TEXT BOOK OF CHEMISTRY best be used only to get a jarful of the gas, mixed with excess of one or other of the ingredients. A stream of the pure gas is easily made by moistening red phosphorus with water, and allowing bromine to fall drop by drop into the paste (Fig. 87). To absorb the bromine vapor, carried by the gas, the latter is passed through a U-tube containing dry red phos- phorus mixed with broken glass 01 beads : PBr 3 3Br 2 3H 2 2PBr 3 . 3HBr H 3 P0 3 . FIG. 87. The bromine forms phosphorus tribro- mide, which is immediately decomposed by the water. The phosphorous acid H 3 P0 3 remains, dissolved in the water, in the flask. The gas can be collected by upward displacement of air. It might seem that a simpler action would be that of sulphuric acid upon a bromide : H 2 S0 4 + KBr -> KHSO 4 + HBr. This action takes place, but hydrogen bromide, being less stable than HC1, which is made upon this plan, is oxidized by the con- centrated sulphuric acid, so that, although some of the gas escapes oxidation, it is mixed with much free bromine and sulphur dioxide : H 2 S0 4 + 2HBr -> S0 2 + 2H 2 + Br 2 . This action, indeed, en- ables us to recognize a bromide, by the color of the bromine vapor and the fuming of the hydrogen bromide produced. A more delicate test is mentioned below. Properties of Hydrogen Bromide. The gas, like hydrogen chloride, is colorless t and has an irritating effect when breathed. It is extremely soluble in water, and fumes in moist air, giving a fog of HBr dissolved in water. IODINE I 2 327 Chemically, hydrogen bromide is stable, though not so much so as hydrogen chloride. Its aqueous solution is an active add. As such, it gives double decomposition with bases and salts. In these actions the bromides, if insoluble, are precipitated. Thus, with a salt of silver, we get a cream-colored precipitate: AgN0 3 + HBr - AgBr J + HNO 3 of silver bromide (used on photographic plates and films). Chlorine-water, added to a solution of any bromide, displaces the bromine, which may be recognized by its brown color (test for a bromide) : Cl 2 + 2Br--2Cl- + Br 2 . A few drops of carbon disulphide, shaken with the mixture, will settle to the bottom, carrying the brown bromine with it in a more concentrated, easily recognizable form. Uses of Compounds of Bromine. Bromine is manu- factured in large amounts in Germany and in the United States. It is employed to make potassium bromide, and other bromides. These are utilized in medicine, and to precipitate silver bromide in the manufacture of photographic films and plates. IODINE I 2 Sources. Iodine was formerly all obtained from seaweed (kelp), certain species of which use the traces of organic com- pounds of iodine in sea water as part of their food. The dried seaweed is carbonized in retorts, and sodium iodide remains in the residue, along with much sodium carbonate and carbon. In an improved process the iodine compounds are dissolved out of the kelp, and from the latter a sort of gelatin, named algin, is extracted. The greater part of our supply of iodine is at present obtained from potassium iodate KlOa, which forms about 0.2 per cent of crude Chile saltpeter. 328 TEXT BOOK OF CHEMISTRY Preparation. The processes for obtaining iodine from the sodium iodide are precisely the same as those for bromine. In France chlorine is used to displace the iodine : Cl 2 + 2NaI-2NaCl + I 2 |. The precipitate of iodine is pressed free from the solution. In Great Britain the iodide is heated with manganese dioxide and sulphuric acid and heated: 3H 2 S0 4 + MnO 2 + 2NaI -> MnS0 4 + 2NaHS0 4 + 2H 2 + I 2 | Iodine vapor condenses upon a cold surface, not to the liquid, but directly to the solid, crystalline form. Distillation which gives a solid product is called sublimation. The crude iodine is purified by repetition of this process. Physical Properties. Iodine forms black, shining crystals. The vapor, visible even at the ordinary temperature, is violet in color (hence the name of the substance, from Greek, meaning like a violet) and has a density corresponding to the formula I 2 . The substance is very slightly soluble in water. It dissolves, however, easily in carbon disulphide (violet solution), in alcohol or potas- sium iodide solution (brown solution), and even in starch, upon which a trace of it confers a strong blue color (test for free iodine, p. 220, see also p. 219). These colors are shown only by free iodine the iodides are colorless and do not act upon these solvents. Chemical Properties. Iodine unites very slowly and in- completely with hydrogen, giving hydrogen iodide. It combines readily with phosphorus (Pis) and with many of the metals, giving iodides. An aqueous solution of potassium iodide, hydrogen iodide, or any other iodide, has likewise the power to take up large quantities of iodine. Here the formation of definite compounds (such as, KI -f I 2 < KI 3 ), by a reversible action, accounts for the amount of iodine taken up. IODINE I 2 329 Hydrogen Iodide, Preparation. The gas is prepared by the process used for hydrogen bromide. Red phosphorus and iodine are mixed, and water is allowed to fall drop by drop upon the mass (Fig. 25, p. 52) : 2P + 3I 2 ->2PI 3 . PI 3 + 3H 2 O -> 3HI + H 3 P0 3 . The gas is very dense (HI = 1 + 127 = 128 g. per 22.4 -liters, against 28.95 g. for air) and can be collected by upward displace- ment of air. The action of sulphuric acid upon an iodide does not give pure hydrogen iodide, because most, though not all, of the latter is oxidized by the sulphuric acid. Hydrogen iodide, being much less stable than even hydrogen bromide, is a more active reducing agent, and reduces the sulphuric acid to hydrogen sulphide. The odor of this gas is therefore very conspicuous when an iodide is moistened with sulphuric acid: Nal + H 2 S0 4 - NaHS0 4 + HI. H 2 S0 4 + SHI -> H 2 S + 4H 2 + 4I 2 . The violet vapor of iodine becomes visible if the test-tube is warmed. A rough test for an iodide is afforded by this action. Still another method of making hydrogen iodide is frequently employed when a solution of the gas in water is required, and not the gas itself. Powdered iodine is suspended in water, and hydro- gen sulphide gas (q.v.) is introduced through a tube in a continuous stream. The iodine dissolves slowly in the water, I 2 (solid) +2. 1 2 (dslvd.), and acts upon the hydrogen sulphide, which likewise dis- solves, H 2 S (gas) ^ H 2 S (dslvd.). Sulphur separates in a fine powder, S (dslvd.) + S (solid), and hydrogen iodide is formed in accordance with the equation: This action takes place, however, only in presence of water, al- though the water does not appear in the equation. The solution 330 TEXT BOOK OF CHEMISTRY is freed from the deposit of sulphur by filtration, and may be con- centrated to 57 per cent of hydriodic acid by distilling off the water. Properties. Hydrogen iodide is exceedingly soluble in water, and fumes strongly in moist air, giving a fog of HI solution. It is colorless. The aqueous solution is an active acid. The iodide-ion I~, which it contains, gives precipitates of insoluble iodides, like the yellow silver iodide Agl : Chlorine-water or bromine-water, added to a solution of this or any other iodide, displaces the iodine : The free iodine, even if present in minute amounts, may be recog- nized by shaking the liquid with a few drops of carbon disulphide. The iodine gives a violet solution in the latter. A still more delicate test is the addition of a drop of very thin starch paste, which gives a deep-blue tint with free iodine. Filter paper dipped in starch paste and dried can also be used, by touching it with a drop of the solution containing the free iodine. Uses of Iodine and Its Compounds. The alcoholic solution (tincture of iodine), painted over the skin, reduces swellings and inflammation. lodoform CHI 3 is a solid used for similar purposes. lodothyrin is an organic compound found in the human thyroid gland, as well as that of other animals. An extract of sheeps' thyroids is administered with remarkable success in cases of degeneration, associated with abnormally small natural develop- ment of the gland (cretinism). Potassium iodide is also used in medicine, to cause absorption of blood-clots and effusions of blood, for example in the eye. Silver iodide is contained in the coating on photographic plates and films. FLUORINE F 2 331 FLUORINE F 2 Compounds of fluorine are found in large quantities as minerals, but the compounds are so stable that the element is very difficult to liberate. The natural compounds, however, have many im- portant uses. Occurrence. Calcium fluoride CaF 2 (fluorite) occurs in nature in beautiful cubical crystals. Cryolite A1F 3 , 3NaF is used in the modern manufacture of aluminum (p. 422). Apatite Ca 5 (PO 4 ) 3 F is a common constituent of rocks and soils. When slowly decomposed, by weathering, it furnishes soluble phosphates. These are absorbed by plants, for which they are a necessary food. Preparation of Fluorine. The element is obtained by electrolysis of potassium-hydro- gen fluoride KHF 2 dissolved in liquefied hydro- gen fluoride. The electrodes are made of an alloy of platinum and iridium, with which fluorine has little tendency to combine. The vessel is a U-tube of copper (Fig. 88) and, to prevent vaporization of the hydrogen fluoride (b.-p. 19.4), is kept at 23 to 40 during the operation. Hydrogen is liberated at one pole and fluorine at the other. After long and fruitless trial of other methods by many chemists, this one was discovered by Moissan (1886). Properties. Fluorine is a yellow gas, with a density greater than that of air (G.M.V. weighs 38 g.). It is the most difficult of the halogens to liquefy (b.-p. 187). Fluorine is the most active of the non-metals. It combines with all the metals, but most slowly with platinum and with gold. In FIG. 88. 332 TEXT BOOK OF CHEMISTRY the preparation of the gas the copper is protected from serious attack by the layer of fluoride first formed. Fluorine combines also with hydrogen in the cold and, unlike chlorine, without the assistance of light. It combines with most of the non-metals, but not with oxygen, chlorine, nitrogen, or the indifferent gases of the atmosphere. With water (vapor or liquid) fluorine interacts, giving a deep- blue gas, ozone, and hydrogen fluoride H 2 F 2 : Skeleton: F 2 + H 2 0-+ H 2 F 2 + 3 . Balanced: 3F 2 + 3H 2 -> 3H 2 F 2 + 3 . Hydrogen Fluoride H 2 F 2? Preparation. When pulverized calcium fluoride and concentrated sulphuric acid are placed in a retort of platinum or lead and the mixture is warmed, hydrogen fluoride passes over. The vapor is usually led into water, in which it is very soluble (hydrofluoric acid) : CaF 2 + H 2 S0 4 -* CaS0 4 + H 2 F 2 . The acid is kept in bottles of paraffin or rubber, as glass interacts with it rapidly. Physical Properties. The vapor of hydrogen fluoride can be condensed to a colorless liquid boiling at 19.4. Being very soluble, it fumes strongly in moist air. The vapor density below 40 corresponds with the formula H 2 F 2 , but at higher temperatures gradual dissociation to HF occurs. Chemical Properties. The aqueous solution has all the properties of an acid. The substance has, in addition, the re- markable property of acting upon silica SiO 2 (sand), and silicates, to give silicon tetrafluoride SiF 4 (a gas). Hence it attacks glass, which is a mixture of sodium silicate and calcium silicate : Si0 2 + 2H 2 F 2 -> SiF 4 + 2H 2 O. CaSi0 3 + 3H 2 F 2 -> SiF 4 + CaF 2 + 3H 2 O. THE HALOGENS AS A FAMILY 333 Thus, when glass is covered with melted paraffin to protect the surface, and marks or letters are made by removing the paraffin with a sharp instrument, hydrogen fluoride will decompose the glass at the parts thus exposed (test for fluorine) . In this way the graduation on thermometer stems and lettering on glass are frequently made. The vapor gives rough, easily visible depres- sions, the solution smooth, glossy ones. On account of this property, hydrofluoric acid is used for re- moving adhering sand from castings and for cleaning the outsides of granite and sandstone buildings. THE HALOGENS AS A FAMILY In the second page of this chapter (p. 324) it was shown that the heat of combination of potassium with the halogens was greatest with fluorine, and least with iodine. The affinity of the halogens for all positive elements diminishes in the order F, Gl, Br, I, that is as the atomic weight of the halogen increases. We have learned also that chlorine displaces bromine and iodine from bromides (p. 327) and iodides (p. 328) . In like manner, bromine displaces iodine. It may be noted that oxygen displaces iodine from hydriodic acid, and that all the halogens displace sulphur from sulphides. Thus an order of activity for the non- metals has been established, similar to the activity order for the metals. For the non-metals it is : F, 01, Br, O, I, S. Although the halogens thus show different degrees of activity, they are closely alike in their chemical nature. They are all univalent; * they are strongly negative or non-metallic; their hydrides are all active acids. They differ in two of these points from other non-metals, for oxygen and sulphur are bivalent (and the latter sexivalent also), and the hydrides of oxygen (H 2 O) and of sulphur (H 2 S) are very feeble acids. * They are so when in binary combination with a metal: KBr, KI, ZnBr 2 , etc. In oxygen compounds they often show a higher valence, as in KClOj and HC1O 3 (chloric acid). 334 TEXT BOOK OF CHEMISTRY Exercises. 1. Write the formulae of the bromides, iodides, and fluorides of zinc, carbon, and aluminium. 2. How should you distinguish by chemical reactions the chloride, bromide, iodide, and fluoride, (a) of hydrogen, (6) of sodium from one another ? 3. Write the equation for the action of iodine upon a solution of hydrogen sulphide. 4. Why does hydriodic acid, when left in the air, become brown in color? 5. Make a list of all the acids we have encountered, and note which are weak and which strong. 6. How should you make potassium bromide, starting with, (a) potassium and bromine, (b) hydrogen bromide, (c) potassium iodide ? 7. What is the valence of chlorine in chloric acid (p. 333, footnote) ? CHAPTER XXVIII PHOSPHORUS, ARSENIC, ANTIMONY, BISMUTH IN the preceding chapter we learned that the elements can be grouped in families* and we became acquainted with one family, namely, that to which chlorine belongs. We now take up the elements which, with nitrogen, form the nitrogen family. These elements all have two regular valences, being trivalent and quin- quivalent. Nitrogen, phosphorus and arsenic are non-metallic elements, that is, they do not form positive radicals of salts. Antimony is non-metallic, although in its trivalent condition it acts also as a metallic element. Bismuth is metallic. PHOSPHORUS P. Occurrence. Calcium phosphate Ca 3 (PO 4 ) 2 forms about 25 to 27 per cent of the material of the bones and teeth of animals. The same salt occurs in deposits, as a mineral, and is found scat- tered through all fertile soils. Complex organic compounds of phosphorus, such as lecithin, are essential constituents of the muscles, nerves and brains of animals and are found also in plants. The average man's skeleton contains 1400 g. of phosphorus, his muscles 130 g., and his nerves and brain 12 g. Amongst foods, egg-yolks and beans contain an unusually large proportion, nuts, peas, and wheat (entire grain) coming next. Phosphorus was discovered by Brand in 1669, and by Kunkel in 1670, by distilling at a white heat the solid residue from evapo- rated animal matter. They were both searching for the philoso- phers' stone. Scheele in Sweden prepared it from bones in 1771. * All the families are shown in the table opposite the inside of the rear cover of this book. 335 336 TEXT BOOK OF CHEMISTRY The element is used chiefly in the manufacture of matches and, to a small extent, in roach paste and rat poison. Manufacture of Phosphorus. Phos- phorus is now manufactured by mixing natural calcium phosphate with sand (Si0 2 ) and coke, and heating the mixture in an electric furnace (Fig. 89). The mixture is admitted by moving the traps below the hopper, and is carried into the furnace by the worm conveyor. The resistance of the mass between the electrodes causes great development of heat. The actions may be shown by partial equations, which, when added together, give the complete equation: FIG. 89. Ca 3 (P0 4 ) 2 3CaO + 3SiO 2 P 2 O 5 + 5C 3CaO + P 2 5 SCaSiOa (calcium silicate) 2P + SCO Ca 3 (P0 4 ) 2 + 3Si0 2 + 5C -> 3CaSi0 3 i + 2P T- + SCO T The calcium silicate is melted and runs out below as a slag. The phosphorus (vapor) and carbon monoxide (gas) pass off through the opening near the top. The phosphorus vapor is condensed under cold water. White Phosphorus. The product, after purification, is a colorless, transparent waxy solid (sp. gr. 1.83), which melts at 44 and boils at 287. It is insoluble in water, but dissolves in carbon disulphide. It has a strong odor, resembling ozone. White phosphorus oxidizes in the air, giving, when moist, phosphorous acid and phosphoric acid, and emitting a faint light from which the element derives its name (Greek, light bearer) . It catches fire at a low temperature (about 35), and in burning forms a cloud of the solid phosphorus pentoxide P 2 O 5 . It com- PHOSPHORUS P. 337 bines readily, even when cold, with the halogens, and when heated it unites with sulphur and the more active metals. White phosphorus is a very active poison (fatal dose, 0.15 g.). When traces of the vapor are breathed day after day, a disease, frequently shown by workers in match-factories and consisting in ulceration of the bones of the jaw, makes its appearance. There is no longer any excuse for exposing the work people to the risk of contracting this horrible disease, for matches can now be made without the use of free, white phosphorus (see below) . The use of white phosphorus is forbidden by law in Sweden, France, Great Britain, and Switzerland and is penalized by a special tax in the United States. Red Phosphorus. When white phosphorus is heated to 230 to 300 in a tightly closed vessel (air excluded) it changes into red phosphorus. This material is composed of small crystals, of dull red color, and variable specific gravity 2.19 to 2.34. It is insoluble in carbon disulphide, has no odor, and is not poisonous. On dis- tillation the vapor condenses to white phosphorus. This allotropic form of phosphorus is formed from the white variety with liberation of much heat. It thus contains less energy, and is much less active. It burns to form the pentoxide, but has to be heated to about 240 before it will catch fire in the air. It combines also with elements other than oxygen much less readily than does white phosphorus. Manufacture of Matches. These are of two kinds, ordinary matches, which strike on any rough surface, and " safety" matches. Ordinary matches are still made in some countries by dipping the splints of wood in melted paraffin, and then in a paste made of 4 to 7 per cent of white phosphorus, lead dioxide about 50 per cent, water, and dextrin (paste). The head, when dry, is dipped in varnish to exclude air and moisture. WTiere the use of white phosphorus is very properly forbidden, a sulphide PiSa is sub- 338 TEXT BOOK OF CHEMISTRY stituted. When the match is struck, the friction explodes the mixture of phosphorus trisulphide (combustible) and lead diox- ide (or other oxidizing agent), and the resulting heat sets fire to the paraffin and this, in turn, to the wood. Safety matches carry no phosphorus, but only a mixture of substances containing oxygen, such as potassium chlorate or potassium chromate, with a combustible, like antimony trisul- phide, some dextrin and a filling (e.g., chalk). The box is coated with a mixture of red phosphorus, antimony trisulphide, dextrin and filling. The friction converts a trace of the red phosphorus into the white variety, and the latter sets fire to the head. Phosphorus Pentoxide PzO^,. When phosphorus is burned in dry air or oxygen, under a bell jar, the cloud of pentoxide slowly settles as a white powder. If the pentoxide is thrown into cold water, chemical union takes place. The heat developed pro- duces a hissing sound, caused by the formation and condensation of minute bubbles of steam. The solution when boiled, evapo- rated, and held at 30 to 40, gives crystals of orthophosphoric acid H 3 PO 4 (m.-p. 42). P 2 O 5 + 3H 2 O 4 2H 3 PO 4 . On account of its tendency to unite with water, the pentoxide is used for drying gases. Phosphates of Calcium. It will be seen that natural calcium phosphate Ca 3 (P04)2 is the orthophosphate of calcium. It is found in considerable deposits in So. Carolina, Florida, Tennessee, and several western states, and in Algeria and Tunis. Since all plants contain organic compounds of phosphorus, they must find a continuous supply in the soil, otherwise their growth will cease. Hence, when the proportion present is too small, calcium phosphate is frequently added as a fertilizer. Since, however, plants can absorb solutions through their roots, but not PHOSPHORUS P. 339 solids, this insoluble salt affords only an exceedingly dilute phos- phorus diet. Hence, a more soluble compound is to be preferred. This is found in calcium acid-phosphate (" superphosphate") CaH4(P0 4 )2, which is made by heating pulverized natural calcium phosphate with sulphuric acid, containing the requisite propor- tion of water: Ca 3 (P0 4 ) 2 + 2H 2 S0 4 + 4H 2 -> CaH 4 (PO 4 ) 2 + 2CaS0 4 ,2H 2 0. The whole turns into a dry mixture, consisting of the superphos- phate and gypsum (hydrated calcium sulphate). The latter does not interfere with the fertilizing power of the soluble superphos- phate, so the mixture is placed directly in sacks and sold as " super- phosphate of lime." In slaughterhouses the bones, after being deprived of fat and gelatin, give a residue containing much calcium phosphate. This residue is treated with sulphuric acid and made into fer- tilizer. Test for a Phosphate. Most phosphates (and phosphoric acid), when mixed intimately with dry sodium carbonate, char- coal, and magnesium powder, and heated in a narrow tube closed at one end, give a phosphide of the metal (Na 3 P or Mg 3 P 2 ) . When the mass is moistened, the odor of phosphine PH 3 can be recog- nized. Na 3 P + 3H 2 - 3NaOH + PH 3 f . Phosphorus and Nitrogen Compared. Although the simple substances, phosphorus and nitrogen, have little in com- mon, they form compounds of similar composition, which are in many ways alike. Thus we have ammonia NH 3 and phosphine PH 3 , both gases. Then there are the oxides N 2 O 3 and N 2 5 , to which the oxides P 2 O 3 and P 2 C>5 correspond. All these oxides are anhydrides of acids. Both nitrogen and phosphorus are typically non-metallic elements (p. 335), entering into negative radicals. Both elements are trivalent in one series of compounds (NH 3 ,N 2 0.3, 340 TEXT BOOK OF CHEMISTRY PH 3 , P20 3 ) and quinquivalent (N 2 5 , P20 5 ) in another series. For these and other reasons they are classified in the nitrogen family of elements. ARSENIC As Arsenic, the third member of this family, is metallic in appear- ance and in physical properties, although in combination it behaves as a non-metallic element. The metal is used with lead in making small shot, and the oxide As20s in preparing medicines and in- secticides. Preparation. A mineral similar to pyrite FeS2, but contain- ing arsenic in place of half of the sulphur, arsenical pyrites FeSAs is one of the commonest natural forms of arsenic. When this mineral is heated (air excluded), arsenic passes off as vapor and condenses as a crystalline metallic powder: FeSAs - FeS + As | Most natural sulphides (PbS, FeS 2 , SnS 2 , CuFeS 2 , etc.) contain more or less arsenic, which takes the place of a part of the sulphur. When these ores are oxidized in a draft of air (roasted), as one step towards the ultimate extraction of the metal, the metal, sulphur, and arsenic are all converted into oxides. The sulphur dioxide passes off as gas, but the arsenic trioxide As20a, a solid, settles in the flues. By distilling the deposit with carbon, free arsenic is obtained: As 2 O 3 + 3C - 2As t +3CO |. Properties and Uses. The element has a silvery luster, but tarnishes quickly. When it is heated, its vapor reaches a pres- sure of 760 mm. before the melting-point is attained, so that the metal distils without melting. The metal burns in air with a bluish-white flame, giving clouds of white particles of the trioxide As2O 3 (poisonous). In the making of small shot, about 0.5 per cent of arsenic is added to the lead. The latter is then run into a vessel, with a ANTIMONY SB. 341 perforated bottom, placed at the top of the shot tower. The arsenic, like any dissolved substance (p. 112), lowers the freezing- point of the solvent (lead), and delays the solidification of the lead until the drops have assumed perfect spherical form. At the foot of the tower the drops fall into water and are cooled. The arsenic also renders the metal harder than pure lead, and less apt to be deformed during the explosion of the cartridge. Arsenic Trioxide As 2 O 3 . This oxide, formed when arsenic burns, is a white crystalline powder. It is acidic and forms salts with bases. Paris green and Scheele's green are made by dissolv- ing the oxide in boiling water and adding a salt copper. They are green, insoluble compounds used, as insecticides, for spraying plants. On account of their poisonous character, they are no longer employed as ingredients in paints. Arsenic trioxide (white arsenic) is the poison most commonly used for criminal purposes, yet it is one of the easiest poisons to detect in the body after death. It is contained in "Fowler's solution" which is used as a heart tonic. . . r. . Arsenic as a Member of the Nitrogen Family. While free arsenic is as different physically from phosphorus as the latter is from nitrogen, the compounds have much in common. Arsenic forms arsine AsH 3 (corresponding to NH 3 and PH 3 ) and arsenious oxide AsjjOs and arsenic oxide As20 5 , and the acids HsAsO 3 and H 3 As04. In these compounds it is non-metallic, and shows the valences three and five. ANTIMONY SB Preparation and Properties. Antimony is found free in nature. The sulphide, stibnite 80285, is also a well-known mineral. When the latter is melted with iron, ferrous sulphide and free antimony are formed. 3Fe + 80283 - 2Sb + 3FeS. 342 TEXT BOOK OF CHEMISTRY The molten ferrous sulphide (sp. gr. 4.8) floats upon the molten antimony (sp. gr. 6.5), and the products, being mutually insoluble, are easily separated. The metal is brilliantly silvery and non-tarnishing. It is brittle and the black powder obtained by pulverizing it, "antimony black," is rubbed on plaster casts to give them a dull, metallic appearance. When heated, antimony volatilizes and burns in the air with a brilliant white light to form the white trioxide Sb 2 3 . The trisulphide Sb 2 S 3 (orange-colored when precipitated) is used in making matches and fireworks. Alloys Containing Antimony. The metal is chiefly used as an ingredient in alloys. Lead, when solidifying, shrinks and antimony counteracts this tendency. Hence type metal, which, when cold, must fill the mould completely, is made by melting together 15 to 25 per cent of antimony, 10 to 20 per cent tin, and the rest lead. The alloy is also harder than lead, and is less quickly deformed by handling and by use in the printing press. Babbitt's Metal (Sb 3, Zn 69, As 4, Pb 5, Sn 19), and other anti-friction alloys, used in lining bearings, contain antimony along with zinc, copper, and other metals. Molten mixtures of metals (alloys), when solidifying, do not always form a homogene- ous, solid mass. In an anti-friction alloy, what is wanted is a mass, in general soft, but containing hard particles. The latter bear most of the pressure, yet, as the alloy wears, they are pressed into the softer matrix so that a smooth surface is always presented. An alloy which has the opposite composition, that is, which gives a hard mass containing softer particles, develops heat by friction much more rapidly. BISMUTH Bi As the atomic weight increases, the members of this family become more like metallic elements in their chemical properties. Thus bismuth is a true metallic element. Its oxides are basic, BISMUTH BI 343 and its compounds give positive ions Bi +++ and include salts like the carbonate, sulphate, chloride, and phosphate. Metallic Bismuth. The metal occurs free in nature. It is a brittle metal, with a pink metallic luster. It melts at 270 and vaporizes at a high temperature. It does not tarnish. It is used in preparing alloys with very low melting-points. Thus Wood's metal contains bismuth (m.-p. 270) 4 parts, lead (m.-p. 326) 2 parts, tin (m.-p.. 233) 1 part, and cadmium (m.-p. 320) 1 part. As is the case with other solutions, the melting-point is lower than that of any of the components, namely 60. Alloys of this class are used as plugs in sprinkler systems and stops to hold steel fire-doors open. When, in consequence of a fire, the tem- perature rises, the alloy melts, the water exits are opened and the fire-doors swing shut. Safety plugs in steam boilers, made of a similar, but less fusible alloy, melt when, as the result of failure of the safety valve, the steam pressure, and therefore the tempera- ture, exceeds a predetermined value. They behave in the same way when the water is dangerously low and the metal above the water becomes too hot. Compounds of Bismuth. When strongly heated, the metal burns to form a brown trioxide Bi 2 03. This oxide gives salts with acids. Thus, with nitric acid it dissolves to form a solution of bismuth nitrate: Bi 2 3 + 6HNO 3 - 2Bi(N0 3 ) 3 + 3H 2 O. When the solution is evaporated, the nitrate appears in color- less crystals. If the crystals are placed in water, a white, in- soluble basic nitrate is formed: Bi(N0 3 ) 3 + H 2 ? BiONO 3 I + 2HNO 3 . This is used in medicine under the name of bismuth subnitrate, to cure stomach troubles. The action is reversible, so that when 344 TEXT BOOK OF CHEMISTRY excess of nitric acid is added, a clear solution of the trinitrate is formed. Exercises. 1. By what chemical experiments could you recognize red phosphorus? 2. Explain the fact that solid Wood's metal floats upon melted Wood's metal. 3. Why is white phosphorus always kept under water? CHAPTER XXIX SILICON BORON SILICON belongs to the carbon family, being, like carbon, quad- rivalent and non-metallic. Although silicon does not occur free in nature, yet its com- pounds are so plentiful that about 26 per cent of the terrestrial globe is silicon. Instead of naming all the rocks which contain it, such as sandstone, basalt, granite, and so forth, it is easier to say that limestone is the only common rock which is not siliceous. Silicon Si. The element is now manufactured at Niagara Falls and elsewhere, by heating sand (SiO 2 ) with coke in an elec- tric furnace. The process closely resembles that for making carborundum (p. 193), except that less coke is used: 2C + Si0 2 -*> 2CO t +Si. The element, as prepared in this way, is a grey, crystalline material. Silicon Dioxide SiO 2 (Silica) , Physical Properties. Color- less rock-crystal, often showing large hexagonal crystals is pure silicon dioxide, deposited from natural solutions. When im- purities enter into it, smoky quartz, rose quartz (pink), and amethyst (violet) are formed. Often the impurity changes dur- ing the growth of the deposit, and beautifully variegated spe- cimens, like jasper, catseye, and agate are produced. Chalcedony v opal, and flint contain a small amount of water in combination. The nodules (rounded masses) of flint break in splinters, when struck, and our prehistoric ancestors dexterously fashioned their implements and weapons from this material. The solid structure 345 346 TEXT BOOK OF CHEMISTRY of sponges and diatoms is also hydrated silicon dioxide. The tests of the latter organisms constitute infusorial or diatomaceous earth (Tripoli) used in making polishing powders and for removing coloring matters from oils. Chemical Properties. Silicon dioxide (sand) is not acted upon by acids, with the exception of hydrofluoric acid, which gives silicon tetrafluoride SiF4 and water. Si0 2 + 4HF -* SiF 4 T + 2H 2 0. When silicon dioxide is fused with sodium carbonate, carbon dioxide gas is liberated, and sodium silicate Na 2 Si0 3 is formed' Si0 2 + Na 2 C0 3 -> C0 2 | + Na 2 Si0 3 . The resulting salt is very soluble in water and a strong solu- tion is sold under the name of water-glass or soluble glass. This material is used as a filler in cheap soaps, as an ingredient in artificial stone, a coating to render wood or cloth fireproof, and a cement for uniting glass or porcelain. Eggs are preserved by being submerged in a solution of this salt (1 part sol. of sp. gr. 1.35 : 9 parts Aq). When an acid is added to sodium silicate solution, silicic acid H 4 SiO4 (gelatinous) is precipitated. If this acid is heated, silica SiO 2 remains as a powder. Glass. Calcium carbonate (limestone) interacts at a high temperature with sand in the same way as does sodium car- bonate: Si0 2 + CaC0 3 -> C0 2 | + CaSi0 3 T giving calcium silicate. Now sodium silicate, when alone, is soluble in water. Calcium silicate is insoluble, but forms a brittle, crystalline mass. By using both sodium and calcium carbonates, and employing a larger proportion of sand than that shown in the equation, a material is obtained which has SILICON BORON 347 the qualities required in glass. When cooled, the molten mass becomes viscous and finally, for all practical purposes, solid. Yet it does not crystallize it is amorphous. It is also prac- tically insoluble in water. By pouring the viscous material into moulds, or stamping it with dies, articles of pressed glass are obtained. Bottles are blown, by taking up a sufficient mass of the hot, thick liquid on the end of an iron tube, inserting it in a mould, and blowing until the outline of the mould is filled. Window glass is made by blowing an immense, elongated bubble (6 by 1J ft.), ripping it while still hot and soft, and flattening it out. Plate glass for windows and mirrors is manufactured by pouring out the material upon a east-iron table, with a raised rim, and passing a large, heated iron roller over it. The plate is subsequently ground flat on both sides and polished with rouge (Fe20a). Soda-calcium glass is called soft glass, because it is easily softened by heating. When potassium carbonate is substituted for sodium carbonate, a less fusible substance, used in making some chemical apparatus, and called hard glass, is obtained. When lead oxide is employed in place of the limestone, a potassium-lead silicate K 2 SiOs, PbSi0 3 ,xSi02 is formed which, on being cooled, gives flint glass. This glass has a higher density and greater brilliancy than soft glass, and is used in making vessels of cut glass and lamp chimneys. The cutting is done with a revolving grinding wheel. When glass is allowed to cool quickly, the product is very brittle and apt to crumble to pieces on receiving a shock or scratch. Glassware is therefore all annealed, by being passed on a slowly moving frame through a long furnace, which is very hot at the entrance and much cooler at the exit. During the passage, the molecules of the material adjust themselves to a stable arrange- ment. Colored glass is made by adding oxides of metals which, with the silica, give colored silicates. Oxide of chromium gives green 348 TEXT BOOK OF CHEMISTRY silicates, oxides of copper and of cobalt blue silicates, and oxide of manganese violet. Gold oxide is reduced to the metal, which goes into colloidal solution and gives ruby glass. Milky glass is made by adding calcium fluoride, or stannic oxide. The green color of bottle glass is due to iron (ferrous silicate) derived from impure sand or limestone. The rough surface of ground glass is produced with a sand blast. For engraved glass, the surface is covered with paraffin to pro- tect it from the sand blast, and only the pattern is left exposed. In granite iron ware the surface is covered with a thin layer of easily fusible glass (enamel, see borax). Pure quartz can be melted in the oxy-hydrogen blowpipe, and recently chemical apparatus (silica ware) has been made out of it. It has the advantage of being less soluble than glass, and of not breaking even when it is heated white hot and quenched in cold water. Glass breaks when chilled, because the parts first cooled shrink considerably and a great strain is produced. Quartz suffers very little change in volume with change in temperature, and so unequal cooling causes almost no strain. BORON B. The element boron resembles silicon and graphite in appearance. It has no applications. Boric Acid HBO. This acid is contained in the steam which issues from the ground in certain parts of Tuscany. It is caught in water, placed in basins built over the "soffioni," and separated by evaporation. Much of it is also made from borax. Boric acid crystallizes in white, slippery scales. It dissolves somewhat in water (4 : 100 at 18), and the saturated solution, mixed with an equal volume of water, is used as an ej'e-wash, for example when the eyes have been irritated by dust. It is a mild antiseptic, and preserves foods by preventing the develop- ment of moulds and bacteria. BORON B. 349 When heated, it gives tetraboric acid: Skeleton: H 3 BO 3 - H 2 B 4 O 7 + H 2 O Balanced: 4HsB0 3 -> H 2 B 4 O 7 + 5H 2 O and eventually boric anhydride B 2 Oa. Borax. This salt, when crystallized from solution, combines with water to form a hydrate. As sold, it is therefore the deca- hydrate of sodium tetraborate Na 2 B 4 O7, 10H 2 O. It is made by adding calcium borate, found in California, to sodium carbonate solution. The precipitate of calcium car- bonate is separated by nitration, and the solution is concentrated until crystals appear upon cooling. It is a white crystalline salt. It is added to the glass, used for enamelling and glazing, to make it more fusible and easier to spread in a thin layer. It is a preservative. Since it contains but a small proportion of the metallic oxide (Na^O, 2B 2 O 3 ), it combines with other metallic oxides when fused with them. For this reason the powdered salt is sometimes sprinkled on tarnished metallic surfaces which are to be soldered or brazed. The heat, of the bolt or blowpipe melts the borax, and the latter removes the oxides and permits perfect running of the solder over the surface. The borates thus formed are often colored, and the colors afford a means of recognizing the metallic compound which produced them. In chemical analyses a bead of borax, produced by fusion on a platinum wire, is heated with a particle of the un- known compound and its color then examined. The colors are similar to those already described under colored glass (p. 347). When hydrochloric or nitric acid is added to a hot, strong solu- tion of borax, much of the boric acid crystallizes out when the mixture cools: Na^B 4 7 + 2HC1 + 5H 2 O -> 2NaCl + 4H 3 BO 3 I . A borate can be recognized by adding hydrochloric acid to the solution, dipping turmeric paper in the liquid, and drying it. 350 TEXT BOOK OF CHEMISTRY The color is reddish brown and, when touched with caustic soda solution, becomes greenish black. Exercises. 1. Make an equation for the preparation of, (a) lead silicate PbSi0 3 by fusion of litharge PbO and sand, (b) potassium silicate K 2 SiO 3 . 2. The inside surface of the bottle of sodium hydroxide solu- tion becomes etched and dull. To what is this due? 3. What is the valence of boron? 4. Why is not all of the boric acid deposited from a hot solu- tion containing it? 5. Write an equation for the effect of heat upon borax. 6. Why does the addition of borax render a glass more easily fusible? CHAPTER XXX COMPOUNDS OF SODIUM AND POTASSIUM WE have already considered sodium and sodium hydroxide (Chap. XIII). In this chapter we take up the other important compounds of sodium and their uses, and we devote some space also to potassium and its more useful compounds. In general, we shall find that these metals and their correspond- ing compounds are very much alike in properties. The chief differences are that the sodium compounds are usually cheaper, and that, on account of the difference in the atomic weights of the two elements (sodium 23, potassium 39), smaller weights of the sodium compounds suffice for a given use involving chemical in- teraction. For these reasons the sodium compounds find, in most cases, more applications. Sodium and potassium are univalent, and both are very active as metallic elements. Their hydroxides being strongly alkaline, the elements are often called the metals of the alkalies. SODIUM NA Sodium derives its symbol Na from its German name, natrium. All compounds of sodium, when heated with a Bunsen burner, confer a strong yellow tint upon the flame. Occurrence. Sea water contains about 2.5 per cent of sodium chloride NaCl. The same compound is found in extensive de- posits (rock salt) at Stassfurt and Reichenhall in Germany, in Cheshire (England), and in Michigan, Ohio, Kansas and else- where in the United States. Sodium nitrate (Chile saltpeter) occurs chiefly in an extensive bed on the border between Chile 351 352 TEXT BOOK OF CHEMISTRY and Peru, from which almost our entire supply is at present obtained. The preparation and properties of the metal are described else- where (p. 159). Sodium Chloride NaCl. This compound is obtained by allowing water to flow into the deposits of rock salt and then pumping the solution to the surface and evaporating the water. A certain amount is made, by evaporation, from sea water. The solution is first concentrated, either by solar evaporation in shallow pools or, in northern latitudes, by allowing such pools partially to freeze and then removing the ice. The salt crystallizes in white cubes. It is a necessary article of diet. For example, it furnishes the hydrochloric acid required in digestion (p. 127) . Its chief application, however, is in the manu- facture of other compounds of sodium. Sodium Bicarbonate NaHCO ; >, Manufacture. -This salt is manufactured by the interaction of sodium chloride and am- monium-hydrogen carbonate in the Solvay or ammonia-soda proc- ess. In the very concentrated solutions which are used, a great part of the sodium-hydrogen carbonate is precipitated: NaCl + NH 4 HCO 3 + NaHC0 3 j + NH 4 C1. (1) In practice salt is dissolved in water and the solution is satu- rated with ammonia gas. The mixture is placed in an iron tower filled with perforated shelves. Carbon dioxide, made by heating limestone in special kilns, is forced in at the bottom. The perforations split up the gas into small bubbles, and facilitate its solution to form carbonic acid H 2 C03. With the ammonium hydroxide NH 4 OH in the liquid it gives ammonium-hydrogen carbonate : NH 4 OH + H 2 C0 3 -* NH 4 HC0 3 + H 2 0. (2) SODIUM Na 353 The product interacts as in equation (1). The sodium bicar- bonate is precipitated and is freed from the liqucr in filter-presses. The ammonia is recovered for use by treating the residual liquor, containing ammonium chloride, with the quicklime CaO from the kilns. The quicklime, with the water, gives slaked lime Ca(OH) 2 and the latter liberates the ammonia (p. 257) : 2NH4C1 + Ca(OH) 2 -> CaCl 2 + 2NH4OH -> 2NH 3 t V2H 2 O. The waste liquor contains calcium chloride, and a considerable proportion of unused sodium chloride, along with an amount of sodium bicarbonate required to saturate it which is not recoverable. Properties of Sodium Bicarbonate NaHCO 3 . This salt is a fine, white, not obviously crystalline powder, which is only mod- erately soluble in water. It is commonly known as baking soda. It decomposes slowly in an open vessel, even when cold. When heated, it rapidly gives off carbon dioxide and water, and leaves sodium carbonate: Skeleton: NaHCO 3 = NaaCO-, + H 2 CO 3 <=* H 2 O + CO 2 Balanced: 2NaHCO 3 + Na^COs + H 2 O + CO 2 (1) Like all carbonates, when treated with an acid, it gives carbonic acid and this, in turn, gives water and carbon dioxide. NaHC0 3 + HC1 <=> NaCl + H 2 CO 3 <= H 2 + C0 2 T - (2) This properly leads to its use in fire extinguishers. Baking Powders. The purpose of the powder is to generate carbon dioxide in the dough. The bubbles of the gas are re- tained by the sticky gluten of the flour. They expand when the dough is baked, and give to it the open texture which, when the bread is eaten, facilitates access of the saliva to every particle. Baking soda, if used alone, will give off, when heated, half the carbon dioxide it contains (equation (1)). The sodium car- bonate which remains in the bread, however, has an acrid taste. 354 TEXT BOOK OF CHEMISTRY By its action on the gluten in the flour, it gives also a yellow color and an unpleasant odor. Finally, the carbonate of soda tends to neutralize the gastric juice (acid) of the stomach and so to inter- fere with digestion. To obviate these difficulties sour milk (containing lactic acid) is sometimes used in making the dough. Occasionally vinegar (p. 226) is added. Most frequently a baking powder, containing an acid substance along with the soda, is employed. The acid sub- stances contained in baking powders are alum, acid-phosphate of sodium, or potassium-hydrogen tartrate. The last, known commonly as cream of tartar HKC4H 4 Oe, is considered the most satisfactory: HKC 4 H 4 6 + NaHC0 3 -> NaKCJBWe + H 2 C0 3 -> H 2 + C0 2 . The advantages of cream of tartar are, (1) that it is somewhat insoluble, and so, even when wet, it does not act upon the soda until ample time has been allowed for complete mixing with the dough, and (2) that the sodium-potassium tartrate (Rochelle salt) which remains is not known to be in the least harmful. It is important that the soda and cream of tartar should be used in the correct proportions, which can be calculated from the equation. In commercial baking powders a little corn-flour is added, to keep the particles of the other compounds apart and prevent that gradual interaction which otherwise would be bound to occur. Bakers 9 Bread. The " raising" of bakers' bread is effected by adding yeast. The batch is "set" in a warm place for some hours to permit the yeast to propagate and to act upon the sugar in the flour. In this action, as we have seen (p. 224), carbon dioxide and alcohol are produced. A little sugar, molasses, or malt extract is added to the dough, to afford a larger supply of the sugar re- quired for the production of the carbon dioxide. The whites of eggs " raise" cake, without the presence of any SODIUM Na 355 soda, because of the expansion under heat of the bubbles of air entangled with the albumen when the eggs are "whipped." Sodium Carbonate Na^COs. This salt is manufactured by heating sodium bicarbonate. It is made also by the Le Blanc process. Sodium chloride is treated at a red heat with an equiv- alent amount of sulphuric acid, giving sodium sulphate NagSOi. The latter is roasted with powdered coal and lii estone. The coal reduces the sulphate to sodium sulphide N^S, and the latter interacts with the limestone giving sodium carbonate and calcium sulphide CaS. The carbonate is separated by solution in water. The solution, when concentrated, gives crystals of the decahy- drate, washing soda NasCOa, 10H 2 O. The latter, when heated, leaves " soda-ash" or " calcined soda" Na2C03. Sodium carbonate Na^COs is used in making sodium hydroxide (p. 160) and glass (p. 346), and to soften water (see p. 400). As about two-thirds of washing soda is water, it is cheaper to ship the anhydrous form, except when the hydrate is wanted, as for wash- ing. Properties of Sodium Carbonate Solution. The solution is not neutral to litmus, as we might expect, but distinctly alkaline. The explanation is that water is ionized, although very slightly, and interacts with the carbonate to a minute extent: NaaCOs + H 2 < NaOH + NaHC0 3 . The action produces a base and an acid-salt, and in so far is the reverse of neutralization (p. 166) . An acid or acid-salt giving H + , and a base giving OH~, have opposite effects upon litmus. If they were equally ionized, and therefore equally active, they would leave the indicator unchanged. In this instance, however, the acid (carbonic acid) is a very feeble acid and hardly affects litmus, while the base (sodium hydroxide) is a very active base (p. 170). Hence, although there is but little of either substance 356 TEXT BOOK OF CHEMISTRY formed, still the base is easily able to show its presence by turning red (or neutral) litmus blue. Hydrolysis of Salts. This interaction of a salt with water is called hydrolysis (Greek, decomposition by water). All salts are hydrolyzed, at least to a slight extent. This is no more than we should expect, for salts can be formed by an acid and a base giving, by dr able decomposition, a salt and water, and we have already asserted that double decomposition of ionogens are all reversible. The effect is noticed, however, chiefly when the acid and base are of very unequal activity. A salt which, by hydrolysis, gives an active base and a weak acid, furnishes a solution the re- action of which is basic. Thus, sodium chloride solution is neutral: NaCl + H 2 <=> NaOH + HC1 because the products of hydrolysis are both active. But sodium sulphide solution is alkaline : Na^S + H 2 + NaOH + NaHS because sodium-hydrogen sulphide HNaS is a very feeble acid. Conversely, if the salt gives, by hydrolysis, a weak base and an active acid, then the solution is acid in reaction. Thus, the solution of cupric sulphate is acid: CuS0 4 + 2H 2 <=> Cu(OH) 2 + H 2 S0 4 because cupric hydroxide is a feeble base. Sodium Nitrate NaNO 3 . This salt is prepared by recrystal- lizing Chile saltpeter (p. 313). When heated, it gives off oxygen, leaving sodium nitrite NaN0 2 , a compound much used as a source of nitrous acid in the manufacture of organic dyestuffs. Sodium nitrate is at present the chief /source of nitric acid (p. 313). On account of its solubility (p. 434) and its nitrogen content, it is valuable as a fertilizer. It is used also in the manu- facture of cheap grades of gunpowder. Much of it is converted POTASSIUM K 357 into potassium nitrate, as this salt is less hydroscopic, and the gunpowder made from it keeps better. Other Salts of Sodium. Several of them, such as the peroxide (p. 302) \ the silicate (p. 346), and the tetraborate (p. 349), have been described already. Sodium sulphate Na2S04 is used as a substitute for sodium carbonate in making cheap glass. Sodium sulphite N^SOs (p. 276), made by the action of sulphur dioxide on an aqueous solution of sodium hydroxide, is a convenient source of sulphur dioxide and is also used as a preservative. Sodium thiosulphate Na^Os ("hypo") can be ob- tained readily by boiling sodium sulphite solution with sulphur: Na 2 S0 3 + S -> Na&O, and is used in fixing photographs. Sodium cyanide NaNC (preparation, see p. 405) is used in extracting gold from its ores. POTASSIUM K Potassium receives its symbol from the initial of the German word for it, kalium (related to the word alkali). All compounds of potassium confer a violet color upon the Bunsen flame. Occurrence. Certain deposits of rock salt, especially those at Stassfurt, are covered by a layer of other salts, of which potas- sium chloride KC1 (sylvite) forms a large part. The soil in India and other tropical countries contains potassium nitrate, the nitrogen for which is derived from the decomposition of animal refuse. The nitrogen is oxidized by the help of certain bac- teria to nitric acid, and the latter interacts with compounds of potassium in the soil. In the Southern States, during the civil war, the supply of imported niter was cut off and potassium nitrate for gunpowder was cultivated successfully in this way. Felspar KAlSi 3 Os is a common constituent of many rocks, such as granite. When such rock material, contained in the soil, is 358 TEXT BOOK OF CHEMISTRY decomposed by weathering, the felspar gives clay HAlSi0 4 and soluble compounds of potassium. The latter afford an indispen- sable item in the diet of plants. There are immense quantities of felspar available, but the process of weathering is very slow, and in many agricultural regions the soil is therefore deficient in available potash. An economical process for manufacturing com- pounds of potassium from felspar is now being sought. The giant seaweeds of the Pacific coast have been found to take up from the sea water, and finally to yield in their ash, an un- usually large proportion of salts of potassium. Until some suit- able source has been brought into use, however, the world will remain dependent upon the Stassfurt deposits for the greater part of the potassium it requires for fertilizers, gun-powder, and a number of other products. Potassium. The metal is prepared from potassium hydroxide by the same apparatus used in making sodium (p. 158). Its properties are similar to those of sodium, excepting that it is more active (p. 54). It has no uses. Potassium Chloride KC1. This is a white salt, found at Stassfurt. It is the chief source of other compounds of potassium. Potassium Carbonate K 2 CO 3 , Preparation. The water used in scouring wool, leaves, when evaporated, the potassium salts of organic acids. When the residue is roasted, potassium carbonate remains. Again, as we have seen, plants absorb potassium compounds from the soil. Hence, wood ashes con- tain considerable amounts of potassium carbonate, and were indeed originally the only source of this compound. The sugar beet takes up exceptional quantities of potassium and, after all the sugar has been removed from the extract, potassium carbonate is obtained by evaporating the liquid and calcining the residue. Much of the salt is also manufactured direct from potassium chloride. POTASSIUM K 359 Properties of Potassium Carbonate. This salt, being much more soluble than sodium carbonate, is deliquescent. Its aqueous solution is alkaline in reaction. It is used in making soft soap and difficultly fusible glass (p. 347). Potassium Bicarbonate KHCO> When carbon dioxide is led into potassium carbonate solution potassium bicarbonate KHC0 3 is formed: C0 2 + H 2 < H 2 C0 3 + K 2 C0 3 *= 2KHC0 3 . The bicarbonate decomposes easily, especially when warmed, reversing the above action (read the equation backwards). This occurs even in the solution. Hence carbon dioxide can be forced by pressure in large amounts into a warm solution of potassium carbonate, and liberated again by pumping so as to create a vacuum. This plan is used as a means of purifying carbon dioxide (p. 199). The same quantity of potassium carbonate can be used over and over again. Before the nineteenth century, when potassium salts were cheaper than sodium salts, potassium bicarbonate made from wood ashes was used in raising bread. It was known as saleratus (Latin, aerated salt) a name which is now sometimes given to the baking soda which has displaced it. Potassium Hydroxide KOH. This alkali is made like sodium hydroxide (p. 160). Either potassium carbonate is treated with slaked lime, or potassium chloride is electrolyzed in a Castner (p. 160) or in a Townsend or a Briggs cell. The solution is evaporated, and the substance cast in slender sticks. Its properties are much like those of sodium hydroxide. It is used in making soft soap and other compounds of potassium. Potassium Nitrate KNO 3 . The supply of the natural nitrate being insufficient, the salt is made by double decomposition from the Chile saltpeter NaN0 3 NaN0 3 -f KC1 -> KN0 3 + NaCl j . 360 TEXT BOOK OF CHEMISTRY Sodium chloride is not much more soluble in hot water than in cold. The three other salts, however, become very soluble as the temperature rises. Hence, when sodium nitrate and potassium chloride are heated with very little water, they dissolve, sodium chloride is precipitated, and potassium nitrate remains in solution. The mass is filtered quickly through canvas to separate the pre- cipitate, and potassium nitrate crystallizes from the filtrate as it cools. The salt is used in making gunpowder and fireworks. It is employed also in preserving ham and corned beef. Gunpowder. Gunpowder is composed of potassium nitrate (75 per cent), charcoal (15 per cent) and sulphur (10 per cent). The ingredients are moistened with water, and intimately mixed by grinding under the heavy rollers of a mill. The "mill cake" is then broken up and granulated to the requir'ed size. The explosion results largely in the union of the charcoal with the oxygen from the nitrate and of the sulphur with the potassium : 16KN0 3 + 21C + 5S - 5K 2 C0 3 + K 2 S0 4 + 2K 2 S 2 + 13CO 2 + 3CO + 8N 2 . One gram of powder yields 264 c.c. of gas (C0 2 , CO, and N 2 ) measured at and 760 mm., and a much larger volume at the temperature of the explosion. One gram produces about 660 calories of heat. The explosion is due to the suddenness with which the gases are generated and the heat is developed. The smoke is composed of particles of the three solid compounds of potassium and is therefore very slow in dissipating itself. Smoke- less powder produces no solids when it explodes. Gunpowder is still used in mining, and for detonating charges of smokeless powder. Other Compounds of Potassium. Potassium bromide KBr (p. 326), used in medicine to relieve nervousness and to produce sleep, potassium iodide KI (p. 328), used also in medicine, and POTASSIUM K 361 potassium chlorate KClOs (p. 337), employed in making matches and fireworks, have already been mentioned. Potassium sulphate K 2 S04, made by the action of sulphuric acid upon potassium chloride at a red heat, is employed in making alum (p. 426). Exercises. 1. What two methods are used for concentrating a solution? Explain why each is effective. 2. In the Solvay process why is the ammonia dissolved in the salt solution, and not separately in water ? Make the equation for the action of heat on limestone. 3. (a) In what proportions by weight should baking soda and cream of tartar be used in raising bread? (b) What is the objec- tion to using too large a proportion of baking soda? (c) Why must baking powder be kept in a dry, cool place? 4. Why does vinegar liberate carbon dioxide from baking soda? Make the equation for the action. 5. In what way must the moisture in the dough contribute to the raising of bread or cake in the oven? 6. Make the equations for the action of: (a) coal (carbon) on sodium sulphate, (b) sulphuric acid upon sodium carbonate. 7. What is the exact percentage of water in washing soda ? 8. What will be the reactions to litmus of aqueous solutions of: (a) borax, (b) sodium bromide, (c) sodium silicate, (d) sodium peroxide, (e) sodium nitrate? 9. Why does a piece of glass, when strongly heated, confer a yellow color on the Bunsen flame? CHAPTER XXXI THE RECOGNITION OF SUBSTANCES, I. A REVIEW OF THE NON-METALLIC ELEMENTS IN the chapters preceding we have classified our substances under one of the chemical elements they contained. Thus, starch, sugar, and alcohol were put under carbon. Hydrogen sulphide and sulphuric acid were considered as compounds of sulphur. Now this classification is of a theoretical nature. That starch, sugar, and alcohol all contain carbon can not be told by mere inspection. We have to make experiments, and to reason about the results, before we reach this inference. Thus we put our inference as the title of the chapter, and distributed the observa- tions and data through it. There is, however, another way of classifying the facts, which is just the opposite of this one. It is the practical classification. When we obtain a specimen, or when a substance appears in the course of an experiment, we must be able to tell what particular substance it is. If it is a white powder, it may contain almost any of the whole list of elements. It may be any one of several thousands of substances. We can recognize it only by its physical properties (p. 7) and by the physical prop- erties of other substances that we can get from it by interaction with known chemical compounds. We need, therefore, a plan of operation, and this plan must be based upon a classification by physical properties, not by constituents. One benefit of the discussion of such a plan is that it will afford us a review of some of the facts already mentioned, by presenting the same facts from a different view-point, and by showing the uses to which they may be put. To avoid unnecessary repetition, we shall refer frequently to the previous statements of the facts. 362 THE RECOGNITION OF SUBSTANCES, ETC. 363 To understand what is here said, therefore, the reader should look up every reference and reread the statement. 7s the Material a Mixture? The first question that occurs to us is, whether the material is a single substance or a mixture. When it is a mixture, we can often, though not always, very quickly find this out. If it is a solid mass or powder, we examine it with the naked eye and with the help of a lens. If we see two or more kinds of par- ticles, as in granite (p. 4), in a mixture of sand and sugar, or in a piece of rusty iron, the kinds differing in form or color or both, then it probably is a mixture. Whether it appears to the eye to be a mixture, or not, we can next try a solvent, such as water, ether, or carbon tetrachloride. If a part of a small sample refuses to dissolve (e.g. sand), while the rest dissolves (e.g. sugar), we have shown that there are two different sets of physical properties, and therefore two different substances (p. 12) present. As there may be very little of the soluble substance in the mixture, we may not perceive at once that anything has dissolved. So we allow a drop of the liquid to evaporate on a clean watch crystal, and observe whether any residue remains. Again, a magnetized knife blade may pick out certain particles (e.g. iron) and not others. If the material is a liquid, we depend largely on differences in the volatility of different substances to find out whether it is a mixture. If a drop evaporates on a watch crystal, leaving a residue (solid or liquid) which does not evaporate, then it is a mix- ture. If this test fails, because all, or none, evaporates, then we must distil the liquid, with the bulb and most of the stem of a thermometer in the vapor (Fig. 35, p. 68), and note whether the whole comes over at one temperature (single substance, in most cases) or whether the temperature changes as the distillation proceeds (mixture, such as petroleum, p. 209). 364 TEXT BOOK OF CHEMISTRY These are simply a few examples intended to show how, when a practical problem is before us for solution we use physical prop- erties as the basis of our reasoning and classification. The Recognition of a Single Substance. The majority of the substances we have met with are acids, bases, or salts. In identifying a substance of one of these classes, it is convenient to attempt the recognition of the positive and of the negative radicals (or ions) as two almost separate problems. In other words, we investigate one radical at a time. On the other hand, substances which do not belong to any of these classes, such as simple sub- stances (sulphur, carbon, chlorine, etc.), oxides (sulphur dioxide, carbon monoxide, etc.), and many organic substances (e.g. carbon disulphide), are investigated as a single problem. Scope of this Chapter. In discussing the recognition of a single substance we shall, for the present, limit ourselves to the non-metallic elements. We shall consider the elementary sub- stances themselves (sulphur, oxygen, etc.), the oxides of such elements, the few organic compounds described, the non-metallic negative ions, and ammonium-ion. We shall leave out of con- sideration until later (Chap. XLIII) the metallic elements (including As, Sb, Bi, Na and K). We shall also ignore the metallic positive radicals, although one of these (or hydrogen-ion) must inevitably be combined with the negative radical under consideration. External Examination (Solids). The specimen may be a solid, a liquid, or a gas. We should note in which of these STATES it exists under room conditions. What follows applies only to SOLIDS the liquids and gases will be taken up later (pp. 368, 369). Without training in crystallography we can tell little about the crystalline form (p. Ill) of the specimen. But anything con- spicuous, such as needle-shape or cubical formation of the particles, must be noted. THE RECOGNITION OF SUBSTANCES, ETC. 365 The COLOR, if any, is significant. If yellow, the specimen may be sulphur (p. 264), if black, carbon (amorphous or graphite, p. 188), if black and crystalline, with violet vapor, iodine (p. 328). Most substances are colorless. The odor, if any, must be noted. Many salts of ammonium (carbonate, sulphide, etc.) smell of ammonia (p. 358). Some sul- phides (of sodium, potassium, ammonium, etc.) smell of hydrogen sulphide (p. 268) . Some acetates smell of acetic acid and hypochlo- rites of hypochlorous acid (p. 306).* Some chlorides (e.g. PCU and A1C1 3 ), in moist air, give hydrogen chloride, with the odor and fumes characteristic of that substance. Effect of Heating (Solids). A great deal may be learned by heating as much of the specimen as will fill the rounded bottom of a small, dry test-tube (Fig. 90). The substance may fuse. Many substances, such as some chlorides and most hydrates, do so. Continue heating. A sublimate (solid deposit on the cold part of the tube) may appear. Black crystals (from violet vapor) are iodine: white crystals, with the limitations we have set, show the substance to be a salt of ammonium (p. 261). Confirm by smelling, and by adding sodium hydroxide to the original substance (p. 260). (Salts of mercury sublime also.) A reddish-brown liquid condenses, becoming, when cold, a yel- low solid. The substance was sulphur from a sulphide (such as FeSs, p. 455). Water condenses. Hydrates, acid-salts, and some organic com- pounds. Test the water with litmus paper. An acid reaction in- dicates an acid-salt (p. 180), an easily hydrolyzed salt (e.g. FeCl 3 , 6H 2 0) or organic acid. Continue heating, inclining the tube * Many experimental details, essential for the successful performance by a beginner of the tests described in this chapter are here omitted. They will be found in the Author's Intermediate Outline of Elementary Chemistry. 366 TEXT BOOK OF CHEMISTRY mouth downwards, removing condensed water with filter paper until no more comes off, and heat the residue. A gas is given off. The gas may be violet (some iodides) or brown (some nitrates, p. 318, and some bromides). If brown, lower a glass rod dipped in silver nitrate solution into the gas in the tube. Bromine will give a white precipitate, etc. (AgBr). In case of a negative result try the test for nitrates. The gas may be colorless. If it has an odor, it may be ammonia from a salt of ammonium, sulphur dioxide from a bisulphite (p. 275) or from oxidation of a sulphide (p. 269). A stifling odor with fumes may be sulphur trioxide from some sulphates or decomposition products from some organic matters. The gas is colorless and odorless. It may be oxygen (test with long, glowing splinter of wood) from a peroxide, chlorate, per- chlorate, or nitrate (of K or Na). To learn which of these it is, dissolve or suspend a little of the substance in water, add dilute sulphuric acid, and test for hydrogen peroxide (p. 304). In case of a negative result examine the residue (as in p. 367). Nitrous oxide also relights the splinter. In this case the substance was ammonium nitrate (p. 321) and can be vaporized completely, leav- ing no residue. A colorless, odorless gas may be carbon dioxide, coming from a bicarbonate or a carbonate (except of K or Na). Lower a glass rod dipped in lime-water into the gas in the test-tube (white precipitate, CaCOal). The substance carbonizes or chars and gives an odor of smolder- ing wood or burning flesh. The compound is organic. Identify, by properties, starch (p. 220), glucose (p. 220), sucrose (p. 221). Heating may produce no effect. On the other hand, more than one of these effects (e.g. both water and sulphur dioxide from a bisulphite) may be given the same specimens. If heating produces any effect, continue heating until all change ceases, and preserve the residue for use in p. 367. THE RECOGNITION OF SUBSTANCES, ETC. 367 In most cases other distinctive properties will be found on the pages referred to. Effect of Sulphuric Acid on Solids. Fill the rounded bottom (only) of a test-tube with the substance,* add just enough concentrated sulphuric acid to moisten the sample, and warm slightly. A gas (effervescence) which fumes in the breath may be given off. If the gas is brown or yellow, it may be bromine (bleaches litmus paper) mixed with hydrogen bromide from a bromide (p. 326). If violet, with brown deposit, accompanied by an odor of hydrogen sulphide, it is iodine from an iodide (p. 329). If the gas fumes, but is colorless, it may be hydrogen chloride from a chloride (add manganese dioxide to get chlorine, p. 135). It may be hydrogen fluoride from a fluoride (a moistened glass rod acquires white precipitate of silicic acid. It may be nitric acid from a nitrate (when cold, add to a little water and copper, p. 317). The gas does not fume. If yellow it may be chlorine dioxide from a chlorate or chlorine from bleaching powder (p. 307) or nitrous anhydride from a nitrite (p. 320). The gas does not fume and is colorless. An odor of sulphur dioxide indicates a sulphite (p. 275) or sulphide (p. 268). An odor of hydrogen sulphide indicates a sulphide. If the gas is odorless, it may be carbon monoxide (burns, leaving carbon dioxide) from a formate, or oxygen from some oxides or a peroxide, or carbon dioxide from a carbonate. No gas evolved indicates a silicate (p. 346), sulphate (p. 286), phosphate (p. 338), or a basic oxide. Sulphuric Acid on the Residue from p. 366. If the sub- stance gave off oxygen when heated alone (p. 356), add a drop * In case the substance gave oxygen in p. 366, use only a few particles, at first, to find out whether it is a chlorate (C102 is explosive) . 368 TEXT BOOK OF CHEMISTRY or two of concentrated sulphuric acid. If the specimen now gives a yellow gas (nitrous anhydride), the original substance was a nitrate (of K or Na), from which the nitrite was formed by heating (p. 356). If it gives a colorless, fuming gas (HC1), the original substance was a chlorate or per chlorate (p. 30). Examination of a Liquid. Test the specimen with litmus paper. A marked acid reaction may be due to an add, such as concentrated or dilute sulphuric acid (p. 285), concentrated or dilute nitric acid (p. 316), concentrated or dilute hydrochloric acid (p. 163), hydrobromic acid (p. 327), hydriodic acid (p. 329), phosphoric acid (p. 338), sulphurous acid (p. 276), or an organic acid; also an acid-salt (p. 180), or a hydrolyzed salt (p. 343). If it bleaches litmus paper, it may be chlorine-water or bromine- water (odor). If it is markedly alkaline in reaction, it may be a solution of a base (NaOH, KOH, NI^OH, etc.) or a hydrolyzed salt (p. 355). Note the odor. Ammonium hydroxide, hydrogen sulphide solu- tion, sulphurous acid, concentrated nitric acid and concentrated halide acids all have odors. Alcohol, acetic acid, carbon disul- phide, carbon tetrachloride, and hydrocarbons (e.g. gasoline) have odors easily distinguished from those of the foregoing. Evaporate a few drops to dryness on a watch crystal. A solid (or liquid) residue shows that the original substance was a solution in water (or possibly alcohol or some other solvent, if the vapor has an odor indicating this) . If there is a solid residue, a quantity of it may be obtained by evaporating a larger amount of the liquid, and may then be treated as a solid (pp. 365-366). If the specimen leaves no residue, and is not acid or alkaline but has an odor, it may be one of the volatile organic compounds named above. If it is odorless, it may be a solution of hydrogen peroxide (p. 302) or simply pure water. THE RECOGNITION OF SUBSTANCES, ETC. 369 Examination of a Gas. The gas has a color. A brown gas may be bromine or nitrogen peroxide. The former liberates iodine from potassium iodide solution (p. 330), but not bromine from a bromide (insert rods moistened with solution of an iodide and a bromide). The latter becomes deeper brown on warming (p. 319). A greenish-yellow gas is chlorine. It bleaches, and displaces bromine from a solution of a bromide (p. 324). The gas may become colored (yellow or brown) on admitting air. It is nitric oxide (p. 318). The gas may have a distinctive odor. Sulphur dioxide, hydro- gen sulphide, and ammonia, also ozone, are of this kind. The gas may fume in the breath. The chloride, bromide, and iodide of hydrogen do so. Distinguish by dissolving in little water and adding manganese dioxide. The gas may be combustible. Burning with a blue flame in- dicates hydrogen (vessel bedewed with moisture), or carbon mo- noxide (leaving carbon dioxide, test, p. 203). Burning with a slightly luminous flame indicates methane (p. 208) and a very luminous flame (often depositing carbon) indicates ethylene (p. 212) or acetylene (p. 212). The gas may relight a glowing splinter of wood. This is oxy- gen, ozone (odor), or nitrous oxide (p. 321). The first two, with nitric oxide, give a brown gas (p. 318), the last does not. The gas having none of these properties is nitrogen (p. 246). In most cases other distinctive properties will be found on the pages referred to. Exercises. 1. Look up the references, and give the properties (physical as well as chemical), other than those mentioned in p. 365, by which you should recognize: (a) sulphur, (b) carbon, and (c) iodine. 2. Same question in regard to: (a) ammonia, (b) hydrogen sulphide, and (c) hypochlorous acid. 3. Same question in regard to : (a) an acid salt, (b) a hydrate, 370 TEXT BOOK OF CHEMISTRY (c) nitrogen peroxide, (d) bromine, (e) sulphur trioxide, (f) hy- drogen peroxide, (g) ammonium nitrate, (h) carbon dioxide, (i) starch, (j) glucose, and (k) sucrose. 4. Same question in regard to (p. 369) : (a) a bromide, (b) an iodide, (c) a nitrate, (d) bleaching powder, (e) sulphur dioxide, (f) carbon monoxide, (g) a silicate, (h) a sulphate, and (i) a phos- phate. 5. Same question in regard to (p. 320) a nitrite. 6. Same question in regard to (p. 368) : (a) cone, and (b) dil. sulphuric acid, (c) cone, and (d) dil. nitric acid, (e) cone, and (f) dil. hydrochloric acid, (g) hydrobromic acid, (h) hydriodic acid, (i) phosphoric acid, (j) sulphurous acid. 7. Same question in regard to: (a) chlorine-water, (b) alcohol, (c) acetic acid, (d) carbon disulphide, (e) carbon tetrachloride, (f) a hydrocarbon. 8. Same question in regard to: (a) nitric oxide, (b) hydrogen, (c) methane, (d) ethylene, (e) acetylene, (f) nitrous oxide, (g) ozone, (h) nitrogen. CHAPTER XXXII FATS, SOAPS, AND RELATED COMPOUNDS THUS far we have mentioned only one organic acid, namely acetic acid HC0 2 CH 3 (p. 226), which is manufactured by the oxida- tion of ethyl alcohol. But many such acids are known and, as we might expect, their composition places them in series just like the hydrocarbon series. Organic Acids and their Salts. The general formula of the main (saturated) series is H(C0 2 C n H 2 n+i). Thus: For n = 0, H(C0 2 H) Formic acid. n = 1, H(CO 2 CH 3 ) Acetic acid. n = 3, H(CO 2 C 3 H 7 ) Butyric acid. n = 15, H(CO 2 Ci 5 H 3 i) Palmitic acid. n = 17, H(CO 2 Ci 7 H35) Stearic acid. Those containing less hydrogen are unsaturated : n = 17, H(CO 2 Ci 7 H33) Oleic acid. Formic, acetic and butyric acids are liquids. Palmitic and stearic acids are solid. Oleic acid is a liquid. These are all monobasic acids. Those with large molecular weights are insoluble in water. All the acids dissolve in sodium hydroxide solution, however, giving sodium salts. Thus stearic acid gives sodium stearate: NaOH + H(CO 2 Ci 7 H35) -> H 2 O + Na(CO 2 Ci 7 H 38 ). Some other salts are : NaCO 2 H sodium formate, NaC0 2 CH 3 acetate of sodium, Na(C0 2 Ci 5 H 3 i) palmitate of sodium, Na(C0 2 CnH 33 ) sodium oleate. Those salts are all soluble in water and common 371 372 TEXT BOOK OF CHEMISTRY soap is a mixture of the last two with sodium stearate. On ac- count of their extensive use in the laundry, and for other purposes in the household, it is very important that we should gain a knowledge of the modes of manufacture and of the properties of these salts. To accomplish this we must first learn the nature of fat, from which soap is made. Abbreviated Formulse. In discussing these substances it will sometimes be convenient to abbreviate the formulae, so that a fatty acid is represented by HCO 2 R and a salt by NaC0 2 R or Ca(C0 2 R)2- Here R stands for a hydrocarbon radical or group of atoms, such as C n H 2n +]. Esters. If we mix alcohol and acetic acid, no immediate action takes place. But after some days or weeks, we begin to notice a fragrant odor due to the formation of ethyl acetate C2H 5 (C0 2 CH 3 ): C 2 H 5 OH + H(C0 2 CH 3 ) -> C 2 H 5 (C0 2 CH 3 ) + H 2 O. Addition of a little sulphuric acid and warming hastens the action, and the odor becomes evident in a few minutes. The foregoing action, when stated in an equation, looks like a neutralization (p. 177). But it differs sharply from ordinary neutralization in several respects. The alcohol is not a base in its properties, although in mere formula or composition it resembles one. In aqueous solution it does not conduct electricity or show the other properties of a base (p. 179). True neutralization takes place instantaneously, while the foregoing action, and all like it, proceed very slowly. Thus, although acetic acid is a true acid, it is not here interacting with a base. Similar actions take place between any alcohol and any acid (organic or inorganic). Thus: FATS, SOAPS, AND RELATED COMPOUNDS 373 CsHnOH + H(CO 2 CH 3 ) -> CsHiiCCOjCHj) + H 2 amyl alcohol acetic acid amyl acetate (banana oil) C 3 H 5 (OH) 3 + 3H(CO 2 CH 3 ) - C 3 H 5 (C0 2 CH 3 ) 3 + 3H 2 O glycerine acetic acid glyceryl acetate C 3 H5(OH) 3 H- 3H(C0 2 Ci5H 3 i) > C 3 Ii5(C0 2 Ci5H 3 i) 3 -f- 3H 2 O glycerine palmitic acid glyceryl palmitate (palmitin) The glyceryl radical C 3 H 5 m is trivalent, as the equations show. The actions into which it enters are, therefore, similar in form to those in which an hydroxide of a trivalent metal, such as alu- minium, is acted upon by an acid : A1(OH) 3 + 3HC1 - AlCla + 3H 2 O. The products of such actions are named as if they were salts, amyl acetate, etc., but they do not possess the properties of salts. They are not ionized in aqueous solution. To distinguish them from salts they are called esters R'(C0 2 R). Fats and Oils. Now the fats and oils found in the bodies of animals, or pressed from the seeds of plants, are composed mainly of various esters. Thus, beef suet is a mixture of about three- fourths glyceryl palmitate (palmitin) C 3 H 5 (CO 2 Ci 5 H 3 i) 3 , and glyceryl stearate (stearin) C 3 H 5 (CO 2 Ci 7 H 53 ) 3 , and one-fourth glyceryl oleate (olein) C 3 H 5 (C0 2 Ci 7 H 33 ) 3 . Hog lard contains about 40 per cent of the former and 60 per cent of the latter, and is therefore softer. Butter includes the same esters, with about 14 per cent of water. When butter is dried, the remaining fat contains about 8 per cent of glyceryl butyrate (butyrin) C 3 H 5 (CO 2 C 3 H 7 ) 3 . Olive oil contains 75 per cent of olein. Cotton seed oil is similar in nature, and is used as a substitute for olive oil or for butter in cooking. All the fats and oils contain, however, a small amount of the free adds, a fact of which we shall later learn the consequences (p. 281). 374 TEXT BOOK OF CHEMISTRY Oleomargarine is an artificial butter. It is made by straining melted beef fat and allowing it to stand at 24. Much of the stearin crystallizes out and the remaining liquid (the " oleomar- garine") is pressed out and allowed to solidify. The solid is finally mixed with a little of some oil, to render it softer, and is churned with milk to impart the proper flavor. Although it lacks the butyrin, the product is similar in chemical nature to butter, and is just as nutritious and wholesome. Many flavors, essences and perfumes are mixtures of esters. Thus the flavor of pineapples is due largely to ethyl butyrate The fats and oils obtained from animals and plants are thus composed chiefly of esters. They must not be confused with mineral oils like petroleum, which are mixtures of hydrocarbons. Physical Properties and Uses of the Fats and Oils. The fats and oils are all insoluble in water. Those with large molec- ular weights are also practically insoluble in cold alcohol. They are soluble, however, in ether, carbon disulphide, and carbon tetrachloride. Hence "grease spots" are not removed by water or alcohol, but are removed by rubbing with the other solvents. Benzene also dissolves fats, and hence is used in washing (dry cleaning) clothing made of silk or woolen materials. The fats and oils belonging to the class of esters are nutritious, and a certain proportion of them is included in* a healthy diet (p. 443). Large quantities of the same fats and oils are also used to make soap. Chemical Properties of Fats and Oils. The chief chemical property of the fats and oils, and in fact of all esters, is that each . can be decomposed by water to give the alcohol and acid from which it is derived. Thus, when ethyl acetate is boiled with; water, it is slowly decomposed into ethyl alcohol and acetic acid : C 2 H 5 (C02CH 3 ) + H 2 -> C 2 H 5 OH + HC0 2 CH 3 . FATS, SOAPS, AND RELATED COMPOUNDS 375 In the case of the fats and oils, they are heated in a closed vessel so as to secure a high temperature (about 200) : 3i)3 + 3H 2 -> C 3 H 5 (OH) 3 + 3HCO 3 C 15 H 3 i. palmitin glycerine palmitic acid When the mixture has cooled, the acid, which is insoluble in water, forms a solid cake, while the glycerine is dissolved in the water. When tallow (beef fat) is treated in this way, the solid is a mixture of palmitic, stearic, and oleic acids. The latter, being liquid, is pressed out, and the solid material is used with paraffin in making candles. The glycerine is separated from the water and used in making nitroglycerine and in medicine. When the fat is heated with a solution of caustic soda NaOH, instead of water, the sodium salts of the acids are obtained. These sodium salts are known as soaps (see next paragraph), and the operation is called saponification : + 3NaOH -> C 3 H 5 (OH) 3 + 3Na(CO 2 C 15 H 3 i) palmitin glycerine sodium palmitate (a soap) The sodium palmitate or other soap is soluble in the water. When common salt is added, however, the soap coagulates and separates into a floating layer which solidifies on cooling. Soft soap is made with potassium hydroxide, and is composed of the potassium salts of the organic acids. Manufacture of Soap. Soap is made in large iron caldrons. These contain closed steam coils for heating, and pipes delivering live steam, when needed, for stirring the mass. The fat is mixed with caustic soda solution containing, at first, only about one- fourth of the total amount of the base that the above equation requires. When, after heating and stirring, a uniform mixture has been produced, the rest of the alkali is added gradually and the heating is continued until the reaction appears to be complete. Salt is now dissolved in the mixture and the soap separates as a curd. The curd floats, leaving the "spent lye," containing the 376 TEXT BOOK OF CHEMISTRY salt solution and much of the glycerine, in the lower layer. This process is called salting out. When the process stops at this point, the upper layer is known as curd soap, and may be dipped out and allowed to cool and solidify. Most " Marseilles" soaps are curd soaps made in this way. A large part of the imported "Castile" soaps are of the same kind. Curd soaps contain salt, glycerine, adhering lye, and other impurities. To prepare a purer soap, the spent lye is run off, dilute brine is added to the soap, and the curd and brine are stirred up. When separation has again occurred, the brine is run off and the process repeated. Finally, some water is added, and stream is run in until the curd mixes completely with the water. When the solution stands, it " settles," that is, resolves itself once more into two distinct layers. The upper layer is called settled soap. The washing with brine and temporary dissolving in water remove the impurities, and hence settled soap is the purest variety. The greater part of the soap made in the United States and in Great Britain, and much of that made in other countries, is of this kind. The qualities of soaps are varied by adding "fillers," such as sodium carbonate, borax, or sodium silicate. Soap powders are often made of ground soap mixed with sodium carbonate. Dyes and perfumes are sometimes added to soaps. Air bubbles are mixed with the soap, by beating, to give the floating varieties. Soap for scouring contains fine sand. Transparent soaps are made by dissolving the soap in alcohol, or by the addition of glycerine or sugar. Chemical Properties of Soaps. The soaps, being sodium salts, dissolve in water and have the usual properties of salts. Thus, when an acid is added, double decomposition takes place: Na(CO 2 Ci 5 H 3 i) + HC1 4 NaCl + H(CO 2 Ci 5 H 3 i) | . sodium palmitate palmitic acid FATS, SOAPS, AND RELATED COMPOUNDS 377 The adds, being insoluble, are thrown down as white pre- cipitates. When other salts are added, similar actions occur. Thus with calcium chloride solution, calcium palmitate is formed, and being insoluble, is precipitated:' 2Na(C0 2 C 15 H 3 i) + CaCl 2 -> 2NaCl This section is important in connection with "hardness" in water (p. 400). Colloidal Suspension. We have seen (p. 220) that starch can be suspended in water in such a fine state of division that the liquid is transparent, and runs through a filter without leaving any solid on the paper. Yet this " suspension" lacks some of the characteristic properties of a solution. The simplest proof, that this is a case of imitation solution, and not of true solution, is obtained by examining the liquid with the ultra-microscope. A converging beam of strong light is sent through the liquid horizontally (Fig. 91), and the illuminated place is viewed from above through a micro- scope. When the room is dark, and the only light comes from the horizontal beam, a col- loidal solution shows minute points of light. A true solution such as one of common salt or of sugar remains perfectly dark. The points of light are produced by myriads of FlQ minute suspended particles. Solutions of soap, gelatine, and many dyes, blood serum, and innumerable other liquids contain such suspended particles. The particles of a colloid (like starch), when viewed in this way, show a continual trembling or vibrating movement (Brownian movement), which is more distinct the smaller the particles, and is due to the impacts of the molecules of the solvent. Colloidal solutions of metals (gold, silver, platinum) can be made with 378 TEXT BOOK OF CHEMISTRY particles so fine that they are of almost molecular dimensions, and shoot about, rapidly, rebounding sharply from the molecules they encounter. Properties of Colloidal Suspensions. Suspended colloids differ from dissolved matter in several respects. Typical colloids are not retained by filter paper. The suspended matter does not settle, and it does not diffuse into a layer of the pure solvent. When a solution of a colloid is placed in a " diffusion-shell " (test-tube shaped) of parchment, surrounded by pure water, none of the colloid escapes through the minute pores of the shell. Or- dinary, non-colloidal solutes do escape, more or less rapidly according as the mass of their molecules is small or large (salt rapidly, sugar slowly) . In this way a mixture of colloid and non- colloid (say, starch and salt) can be separated, if the water sur- rounding the shell is replaced by pure water at intervals until all the non-colloid has been removed. This method of separation is called dialysis. Finally, matter in colloidal suspension can be coagulated (or flocculated), and is then precipitated. Colloidal Matter in Soap Solutions. Explanation of Salting Out. Soaps are somewhat hydrolyzed in solution. Letting R stand for the hydrocarbon part of the acid radical: Na(C0 2 R) + H 2 <= H(C0 2 R) + NaOH. The free acid HC0 2 R combines with an undissociated molecule of the original salt NaCO 2 R to form an acid salt NaH(CO 2 R) 2 . This acid salt is a colloidal substance, and exists in colloidal sus- pension in the soap solution, in equilibrium with the ions and molecules of the original salt and the NaOH. The capacity for being coagulated and precipitated, which is characteristic of colloidal matter, is shown by the soap. Most sodium salts will coagulate a soap solution and precipitate the soap JEAN PERRIN French physical chemist. Investigator of the Browman movement, 1870- FATS, SOAPS, AND RELATED COMPOUNDS 379 as a curd. The acid salt NaH(CO2R)2 seems to collect (absorb) and carry down with it the most of the sodium hydroxide. As both of the substances on the right of the equation (above) are thus precipitated, the equilibrium is displaced to the right, and the precipitation becomes complete. This explains * the process of " salting out" which plays so large a part in the manufacture of soap. Causes of the Cleansing Action of Soap. The chief use of soap solution is in removing grease and dirt from yarn, cloth, or clothing, and from woodwork and kitchen utensils. Soap solu- tion has two more or less distinct properties, one of which enables it to remove oil or grease (viscous, insoluble liquids), and the other of which enables it to remove dust and dirt (largely minute, solid particles of carbon soot). The former is its emulsifying power, the latter is probably connected with its nature as a col- loidal suspension. The Emulsifying Action of Soap. How Soap Removes Grease. When an insoluble oil, such as kerosene or lubricating oil, is shaken with water it is divided into minute droplets sepa- rated by water from one another. When the shaking ceases, how- ever, the droplets begin to run together and soon the oil and water have separated once more into two layers of transparent liquid. When very dilute soap solution and oil are shaken to- gether, however, the droplets do not run together, but remain permanently suspended. The mixture is opaque and more or less viscous. Such a permanent mixture of two insoluble liquids is called an emulsion. * "Salting out" is often explained as being due to the increased concentra- tion of sodium-ion causing the normal ion-product of the soap to be exceeded. This theory falls to the ground when we consider the facts, namely, that precipitation by excess of one ion is (1) confined to saturated (or nearly satu- rated) solutions, and (2) leaves the solution saturated still, while salting out works with dilute solutions of soap and even then gives almost complete precipitation. 380 TEXT BOOK OF CHEMISTRY Soap solution, when rubbed on oily or greasy goods, emulsifies the grease, converts it into droplets, surrounded by soap solution and separated from the cloth, and so permits it to be washed off. In mayonnaise dressing, which is a thick, almost solid, per- manent emulsion, the olive oil is emulsified by the colloidal matter of the yolks of the eggs which have been dissolved in the vinegar. Adsorption of Colloidal Matter. When a solution contain- ing colloidal substances, such as many dyestuffs and organic coloring matters, is shaken with finely pulverized charcoal or with infusorial earth, the colloidal substance adheres to the surface of the powder (is adsorbed) and the liquid is consequently decolorized. This principle is used in decolorizing the extracts in sugar-refining (p. 222), and in "bleaching" oils.* When dilute soap solution is shaken with infusorial earth and the mixture is filtered, the clear liquid is found to have been de- prived of the soap. The soap is evidently precipitated (adsorbed) on the surface of the particles of the solid. How Soap Removes Dirt. Clearly there is a tendency to cohesion between the colloid in a soap solution and the particles of a fine powder. When there is much of the powder, and little of the soap in solution, the powder takes the soap out of the solution. When, however, there is much of the colloid in the form of soap solution, and little of the solid, and that very finely divided, the same tendency to adsorption exists, only, in this case, the colloidal particles carry off the powder. In short, the dirt is removed by adsorption into the solution. Possible Objections to the Foregoing Explanation. For- merly soap solution was supposed to remove grease (and soot ?) because of its slight alkaline reaction, due to hydrolysis. This ex- * Non-colloidal substances, such as gases (hydrogen sulphide and ammonia), are also adsorbed by charcoal. FATS, SOAPS, AND RELATED COMPOUNDS 381 planation must be given up because: (1) an alkali so dilute that it exists in equilibrium with the free fatty acid, can not possibly saponify the ester contained in a grease spot. (2) Pure alkali of the same concentration (or stronger) has no more emulsifying power than water. Such an alkaline solution will indeed emulsify an animal or vegetable oil (cod-liver oil, cotton oil, castor oil), but it does so by interacting with the free fatty acid always present in such oil (p. 373) and forming therefrom a soap. Such an alkaline solution does not emulsify kerosene, although soap solution does. The emulsifying agency is always a soap. (3) Very dilute alkali has no more effect upon soot than has water but soap solution takes clean (greaseless) soot instantly into permanent suspension. The power of forming an emulsion depends, theoretically, upon the abnormally low surface tension of dilute soap solution. Very dilute alkali has the same high surface tension as has pure water. Exercises. 1. Which reactions of acids are shown by palmitic, stearic, and oleic acids, and which are not shown ? 2. Make equations for the formation of: (a) ethyl formate; (b) glyceryl formate; (c) ethyl stearate. 3. Make equations for the action of superheated water on: (a) stearin, and (b) olein; and for the action of caustic soda on: (c) stearin, and (d) ethyl acetate. 4. Summarize the facts which show soaps to be salts. 5. (a) Make the equation for the hydrolysis of sodium palmitate. (b) What reaction (acid or alkaline) should soap solution possess, and why (p. 356)? 6. Why does " French dressing" (vinegar, salt, and olive oil) give an emulsion, which is much less durable than mayonnaise dressing ? CHAPTER XXXIII EXPLOSIVES AND PLASTICS ALTHOUGH very different, in their ordinary uses, from the fats, modern explosives are nevertheless esters, and closely connected in chemical nature with these inert substances. Stranger still, the material of the celluloid hair comb is chemically an ester of the explosive type. And strangest of all, imitation silk, as lus- trous as the product of the silk-worm, is made by hydrolyzing esters like guncotton. Nitroglycerine. As already mentioned, the alcohols interact with inorganic acids, as well as with organic ones, to produce esters. A familiar illustration is met with in the manufacture of nitroglycerine (glyceryl trinitrate) by the action of glycerine and nitric acid: C 3 H 5 (OH) 3 + 3HN0 3 - C 3 H 5 (N03)3 + 3H 2 O. To assist in the liberation of the water, the nitric acid is mixed with a dehydrating agent. The glycerine then is added slowly to the cooled reagents. The nitroglycerine is an almost color- less oil which floats to the surface of the acid mixture. It is shaken repeatedly with water, in which it is insoluble, and then with sodium carbonate solution, in order to free it from all traces of the acids. Nitroglycerine explodes violently, often from the slightest shock. It owes this power to the fact that its carbon and hydrogen can combine with the oxygen it contains to form carbon dioxide and water : - 6CO 2 + 5H 2 + 3N 2 ( 382 EXPLOSIVES AND PLASTICS 383 The latter are very stable substances and much heat is liberated in forming them. They are both produced as gases and, at the high temperature of the action, they and the nitrogen tend to occupy a great volume or to exert an enormous pressure in the effort to do so. The explosion is also so sudden, compared with that of gun- powder, that nitroglycerine would shatter the breech of a rifle before the bullet had time to move. It also pulverizes rock, instead of breaking it into fragments of usable size. For these reasons, as well as on account of the danger in handling, and impossibility of safely transporting the substance, it is made into blasting gelatine (see below). The old form of dynamite was made by soaking a porous earth (infusorial earth, kieselguhr) with nitroglycerine. Guncotton or Cellulose Nitrate. Cotton is composed of cellulose (CeHioOs),,, a substance which has chemically much in common with glycerine and the alcohols. When cotton is steeped for half an hour in a cooled mixture of nitric and sulphuric acids, it is converted into cellulose trinitrate or guncotton: C6H 7 2 (OH) 3 + 3HN0 3 -> C 6 H 7 2 (N0 3 )3 + 3H 2 O. The fibers have the same appearance as before, but are crisper to the touch. The guncotton is washed thoroughly with water to remove the acids, which would cause slow decomposition and perhaps accidental explosion. Dried guncotton burns briskly (deflagrates) when set on fire. While wet, it can be moulded and cut without danger. It ex- plodes only when "set off" by a small amount of another ex- plosive. Fulminate of mercury Hg(ONC) 2 , used in percussion caps, is commonly employed. By such means the explosion is brought about in wet guncotton as easily as in dry. In pure form guncotton is used only in torpedoes and sub- marine mines. Like nitroglycerine it explodes too rapidly to be used in fire-arms or for blasting. 384 TEXT BOOK OF CHEMISTRY The fact that a flat rock or plate of steel is shattered, when a charge explodes upon its surface, does not show that the sub- stance "explodes downwards only." It exercises pressure equally in all directions, but with such extreme suddenness that the explosion pulverizes the rock before the inertia of the atmospheric air can be overcome and space can be made for the expansion of the gases produced. Smokeless Powder and Dynamite. The violence of gun- cotton is reduced by compressing it, and still more by dissolving it and allowing the solvent to evaporate. Thus, cordite is made by dissolving guncotton (65 parts), nitroglycerine (30 parts) and vaseline (5 parts) in acetone. The resulting paste is rolled and cut into pieces of different dimensions, according to the rate of explosion desired. When the acetone evaporates, the horny cordite remains. These explosives are smokeless because they differ from gunpowder (p. 360) in yielding no solids when they decompose. Blasting gelatine, giant powder, and other forms of dynamite are made by dissolving guncotton in nitroglycerine. Substances like nitrate of sodium or of ammonium and sawdust or flour are added to adjust the rate of explosion so that, for example, coal may be split up, but not shattered. Celluloid, Collodion, Artificial Silk. A special guncotton, less completely " nitrated" by the nitric acid, when worked be- tween rollers with camphor and a little alcohol, forms a viscous solution. When the alcohol evaporates, a transparent colorless solid, celluloid, remains. Photographic films are made by rolling the dough into sheets. Fillers and dyes can be added to the dough and the latter can be moulded to any form. In this way ivory-like or black combs and brush handles, opaque white knife handles, articles of " artificial amber" and so forth can be made. The same sort of guncotton dissolves in a mixture of alcohol EXPLOSIVES AND PLASTICS 385 and ether, giving a solution called collodion, used in photography and in medicine. When collodion is forced under great pressure through minute holes in a steel die, the threads dry as they issue from the openings and can be wound on spools. The product is treated with an alkali, which saponifies the ester (p. 383). This "denitration" destroys the explosive qualities, leaving a material of the com- position of the original cotton. The product is one of the forms of artificial silk, and is at least as brilliant in appearance as the real article.* It can be dyed to any desired tint. Cellulose Acetate. When cotton is treated with acetic acid (in the form of acetic anhydride) it gives another ester, cellulose acetate. The viscous liquid dries to a tenacious film. On ac- count of its waterproof character and non-conductance of electric- ity, it is now used for insulating electric wires. Artificial horse- hair (e.g. for making women's hats) and bristles for hair brushes are made of it. Other Plastics from Cellulose. Although the chemical reactions involved probably do not fall within the class considered in this chapter, some other applications of cellulose may conven- iently be mentioned here. Cellulose dissolves in hot, concentrated zinc chloride solution. WTien the liquid is pressed through a small orifice into alcohol, the cellulose is precipitated in the form of a thread. By carbon- izing, this is made into filaments for incandescent electric lamps. Cellulose, such as paper (wood pulp), is soluble also in a solu- tion of cupric hydroxide in excess of ammonium hydroxide, and is reprecipitated by dilute sulphuric acid. Paper or cotton goods can be passed through, first one and then the other of these liquids, and so receive a tough, waterproof surface. Artificial silk is made by pressing the solution through dies into the precipitant. * Real silk is a protein (Chap. I, p. 1), and is not chemically related to cellulose. 386 TEXT BOOK OF CHEMISTRY Cotton, when dipped in strong sodium hydroxide solution, and then stretched while drying to prevent the shrinkage which otherwise occurs, acquires a brilliant luster and is used in enor- mous quantities under the name of mercerized cotton (discovered by John Mercer). Finally, mercerized cotton, or wood pulp treated with caustic soda, combines with carbon disulphide to give viscose. Viscose dissolves in water, and decomposes in solution giving a plastic form of cellulose. This can be rolled into transparent sheets, made into caps for bottles, moulded into any form, or pressed through dies into solutions of salts of ammonium to give still another form of artificial silk. Bakelite, Another Plastic. Not chemically related to the preceding plastics is bakelite, prepared from formaldehyde CH 2 and phenol C 6 H 5 OH (carbolic acid). Under suitable heat treat- ment the mixture gradually sets to a solid, hard, infusible, resi- nous mass, which is insoluble in all known solvents. Before it sets, it can be dyed or " filled," and it can be applied as lacquer, or moulded to any form. Switchboards, dolls, ornamental buttons, artificial jewels, phonograph records, billiard balls, and stereotyp- ing matrices are amongst the objects into which it is now fashioned. Exercises. 1. What chemical change would occur after mix- ing nitroglycerine with sodium hydroxide solution? Name the kind of reaction and give the equation. 2. When nitroglycerine explodes, in what relative volumes are steam, carbon dioxide, and nitrogen produced? What principle is used in answering this question ? 3. Make an equation for the decomposition of guncotton, similar to that given for nitroglycerine (p. 383). 4. Make an equation for the denitration of guncotton by an alkali. 5. In view of the statement on pp. 383-385, make a list of the substances which do react with cellulose. CHAPTER XXXIV CALCIUM AND ITS COMPOUNDS UP to this point our attention has been given largely to non- metallic elements and to negative radicals (e.g. N0 3 ) and the compounds related to them (e.g. NO, N2O, NCy. Having now disposed of most of these elements and radicals, we enter upon the systematic study of the metallic elements, and the radicals (mainly positive) and compounds which contain them. We have been prepared for this by a detailed discussion of two typical metallic elements, namely sodium and potassium, and by a more casual examination of some of the compounds of arsenic, antimony, and bismuth. The Calcium Family. Calcium belongs to a family of metallic elements which includes also strontium, barium, and radium. This family resembles that to which sodium and potas- sium belong in so far that the metals are second in activity only to the two last named and that the hydroxides are active bases. The chief differences are that the metals of the present group are bivalent and that all the carbonates and many other single com- pounds are insoluble. Compounds of calcium confer a brick-red color upon the Bunsen flame. Calcium. The metal is made by electrolyzing melted cal- cium chloride in a graphite crucible, which forms the anode. The cathode, to which the negative wire is attached, is a rod of iron, one end of which dips into the fused salt. The calcium, liberated at this point, adheres to the rod. The latter is slowly 387 388 TEXT BOOK OF CHEMISTRY raised, in such a way that the calcium always remains in contact with the liquid. In this fashion a long " cabbage-stalk" of cal- cium is finally produced. The metal is slightly harder than lead and has a silver-white luster. It is the fifth metal in the order of electrical conductivity. It decomposes cold water, liberating hydrogen (p. 50). It is easier and safer to use for ^jg ex p er j men t than sodium, because it sinks to the bottom, does not melt, and reacts rapidly but not violently. It tarnishes rather gradually, but burns brilliantly when heated with a blast lamp to a temperature above its melting-point. Calcium Carbonate CaCO 3 . The carbonate is the com- monest compound of calcium. White marble is a pure variety, composed of crystals compactly wedged together. Limestone does not show much crystalline structure and usually contains clay and other impurities. Chalk is made up of shells of minute marine organ- isms. Shells, coral, and pearls are likewise mainly calcium carbonate. Well-formed crystals (calcite, Iceland spar Fig. 92 and dog-tooth spar Fig. 93) are common. Limestone is used for building and in road-making. Much of it is employed in making quicklime, cement, FlG 93 and glass (p. 346), and as a flux in metallurgical operations (see p. 456). Marble, often variegated by the presence of impurities, is used in building and sculpture. As we have seen (p. 198), calcium carbonate reacts with acids to give carbonic acid CaCO 3 + 2HC1 -> CaCl 2 + H 2 CO 3 - H 2 O + CO 2 t CALCIUM AND ITS COMPOUNDS 389 OHM* """" OlQXIOt When heated, all forms of calcium carbonate give off carbon dioxide, and leave calcium oxide (see below) : CaC0 3 <= CaO + CO 2 . Calcium Oxide CaO, Manufacture. The manufacture of calcium oxide or quicklime (i.e. live lime) is one of the most ancient chemical industries. The limestone is thrown into a kiln lined with brickwork (Fig. 94). The flames and heated gases from the fire pass .through the limestone and the carbon dioxide is liberated and carried off by the draft. When this gas is to be used, as in the Solvay process or in the refin- ing of sugar, coke (smokeless) is chosen as the fuel. When no use is to be made of the escaping gas, coal may be employed. The use of as low a tempera- ture as possible is important. A high temperature causes the im- purities in the limestone (the clay, FIG. 94. etc.) to interact with the quicklime and form fusible silicates, which fill the pores and retard the subsequent action of water in slaking the lime. Calcium carbonate gives a pressure of only 25 mm. of carbon dioxide at 700 and one atmosphere at about 900. The action is reversible (see equation), and if the gas lingers in the kiln, the carbon dioxide recombines with the quicklime as fast as it is liberated unless a temperature above 900 is used. When, however, the gas is continually removed, the backward action is prevented, and a lower temperature suffices to complete the dissociation of the compound. Hence a low temperature and a good draft of air through the kiln are essential. Properties and Uses of Quicklime. Calcium oxide is a white, amorphous material. It melts only in the electric arc. 390 TEXT BOOK OF CHEMISTRY When heated strongly, it glows with an unusually brilliant and white light. The Drummond, or oxy-hydrogen light, more com- monly called the lime light or calcium light, is produced by allow- ing a flame of burning oxygen and hydrogen or illuminating gas to play upon a cylinder of calcium oxide (Fig. 95). The gases are contained in iron cylinders, under pressure, and so the : IFI uffirf^ apparatus can be used for illumination or FlG 95 projection where neither electricity nor a local supply of gas is available. When water is poured upon quicklime, it is at first absorbed, and then enters into combination to form calcium hydroxide (slaked lime): CaO + H 2 -> Ca(OH) 2 . So much heat is given out that the excess of water is converted into steam. The quicklime swells and falls to powder. Quicklime is used most largely in making slaked lime for mortar, and also in the manufacture of bleaching powder. Other uses are mentioned under the hydroxide. Quicklime deteriorates when exposed to the air. It combines both with the moisture and carbon dioxide in the atmosphere and becomes air-slaked. Calcium Hydroxide Ca(OH) 2 . The hydroxide is a white, amorphous powder. It is slightly soluble in water (about 1 : 600 at 18), giving lime-water. The solution has a strong alkaline reaction, however, showing that so much as is dissolved is very largely ionized. Milk of lime, a saturated solution with a large excess of calcium hydroxide suspended in it, is employed in many operations (see e.g. pp. 160, 359). As the dissolved part undergoes chemical change, more goes into solution. Being cheap, it is used whenever an alkali is needed, provided a dilute alkali will serve the purpose. It interacts with acids giving salts of calcium, and shows the other properties of a base (p. 260). CALCIUM AND ITS COMPOUNDS 391 Slaked lime is used in making mortar (see p. 391) and alkalies (p. 160) and in purifying sugar (p. 222). It is employed to remove the hair from hides, before tanning, an action which recalls the solubility of wool (sheep's hair) in an alkali * (p. 1). It finds application, also, in softening water (p. 403) and as whitewash. Mortar. Mortar is made by mixing slaked lime with three or four times its bulk of sand, and making the whole into a paste with water. When the water evaporates, a porous, rather crumbly material remains. This, however, at once begins to harden, owing to the action of the carbon dioxide in the air upon the line: Ca(OH) 2 + CO 2 -> CaC0 3 -f H 2 | . The crystalline calcite (CaCO 3 ) adheres to, and is interlaced with the sand, and gives a rigid, though porous, structure attached firmly to the brick or stone. The pores facilitate the penetration of the air into the deeper parts and thus provide, both for the fresh supplies of carbon dioxide required for the continuance of the action shown in the equation, and for a considerable amount of useful ventilation through walls of the building. Calcium Oxalate CaC 2 O 4 . This salt may be observed under the microscope in the cells of many plants. It appears in the form of needle-shaped or of granular crystals. Since it is the least soluble salt of calcium, its formation by precipitation is used as a test for calcium ions. Theory of Precipitation. The precipitation of calcium oxa- late CaC 2 4 , just referred to, is a typical one and may be used to illustrate the application of ion-product constancy (p. 235) to explaining the phenomenon. The same explanation serves for all precipitations. * Depilatory preparations consist usually of calcium sulphide CaS, and are applied as a paste with water. The sulphide is hydrolyzed, and the cal- cium hydroxide dissolves the hair without injuring the skin. 392 TEXT BOOK OF CHEMISTRY The first thing to be remembered is that the precipitate which we observe, however insoluble its material may be, does not include all of the substance, but only the excess beyond what is required to saturate the water. The liquid surrounding the pre- cipitate is always a saturated solution of the substance precipitated. If it were not so, some of the precipitate would dissolve until the liquid became saturated. Thus, for example, when we add am- monium oxalate solution to calcium chloride solution: the liquid is a saturated solution of calcium oxalate, with the excess of this salt suspended in it as a precipitate. Looking at the matter from this viewpoint, we perceive the ap- plication of the rule of ion-product constancy. In this saturated solution (p. 109) the product of the ion-concentrations, [Ca ++ ] X [204"], is constant. If the original solutions had been so very dilute that, when they were mixed, the product of the concentra- tions of these two ions had not reached the value of this constant, no precipitation would have occurred. As a matter of fact, the ion- product considerably exceeded the requisite value, and hence the salt was thrown down until the balance remaining gave the value in question. The rule for precipitation, then, is as follows : When- ever the product of the concentrations of any two ions in a mixture exceeds the value of the ion-product in a saturated solution of the compound formed by their union, this compound will be precipi- tated. Naturally the substances with small solubilities, and there- fore small ion-product constants, are the ones most frequently formed as precipitates. In the case of calcium oxalate, the molar solubility (see Table) is 0.0 4 43. In so dilute a solution the substance, being a salt (p. 174), must be practically all ionized. Each molecule gives one ion of each kind. The molar concentrations of these ionic substances, Ca++ and CgO*", in the solution, when the solid is CALCIUM AND ITS COMPOUNDS 393 also present, must therefore be (practically) 0.0 4 43, each. The product [Ca++] X [CzO^] is thus equal to 0.0 4 43 X 0.0 4 43 or 0.0 8 185. If in mixing the solutions, exactly equivalent quan- tities were not employed, the values of the two factors will not be equal, but the product will in any case possess this value. Rule for Solution of Substances. The rule for solution of any ionogen follows at once from the foregoing considerations, and may be formulated by changing a few of the words in the rule just given : Whenever the product of the concentrations of any two ions in a mixture is less than the value of the ion-product in a saturated solution of the compound formed by their union, this compound, if present in the solid form, will be dissolved. When applied to the simplest case, this rule means that a substance will dissolve in a liquid not yet saturated with it, but will not dissolve in a liquid already saturated with the same material. The value of the rule lies in its application to the less simple, but equally common cases, such as when an insoluble body is dissolved by interaction with another substance (next section). Applications of the Rule for Solution to the Solution of Insoluble Substances. So long as a substance remains in pure water its solubility is fixed. Thus, with calcium hydroxide, the system comes to equilibrium at 18 when 0.17 g. per 100 c.c. of water (0.02 moles per liter) have gone into solution: Ca(OH) 2 (solid) <=* Ca(OH) 2 (dslvd.) <=> Ca++ -f 20H~. But if an additional reagent which can combine with either one of the ions is added, the concentration of this ion at once becomes less, the actual numerical value of the ion-product therefore begins to diminish, and further solution must take place to restore its value. Thus, if a little of an acid (giving H + ) be added to the solution of calcium hydroxide, the union of OH~ and H+ to form water removes the OH~, and solution of the hydroxide proceeds until the acid is used up. There are now more Ca++ than OH~~ 394 TEXT BOOK OF CHEMISTRY ions present, but the ion-product reaches the same value as be- fore, and then the change ceases. If a further supply of acid is added, the removal of OH~ to form H 2 begins again. With excess of the acid, the only stable OH~ concentration is that which is a factor in the very minute ion-product of war, [OH~] X [H + ], which is 0.0 6 1 X 0.0 6 1, or 0.0i 3 l. Hence, with excess of acid, the calcium hydroxide, which requires in general a much higher con- centration of OH~ than this to precipitate it or to keep it out of solution, finally all dissolves. More specifically, if we assume that the calcium hydroxide is wholly dissociated in so dilute a solution (which is nearly true), each molecule forms one ion of Ca ++ and two ions of OH~. That is, each mole of Ca(OH) 2 gives one mole of Ca ++ and two moles of OH~~. As the saturated solution contains 0.02 moles of the base, the molar concentration (assuming complete dissociation) of [Ca++] is 0.02 and of [OH"] is 0.04. Now, the ion-product is the product of the concentrations of all the ions formed, i.e. Ca" 1 "*, OH~, and OH~. The value of the product is therefore [Ca++] X [OH"] X [OH-] or [Ca++] X [OH~] 2 . That is, 0.02 X 0.04 2 = 0.0 4 32. Note that if the molecule gives two (or three) ions of the same kind, the whole concentration of that ion taken, and is also raised to the second (or third) power. This particular action is a neutralization of an insoluble base. But the other kinds of actions by which insoluble ionogens pass into solution all resemble it closely, and differ only in details. The general outlines of the explanation are the same in every case. We proceed now to apply it to the common phenomenon of the solution of an insoluble salt by an acid. Interaction of Insoluble Salts with Acids, Resulting in Solution of the Salt. Calcium oxalate passes into solution when in contact with acids, especially active acids. Thus, with hydrochloric acid, it gives calcium chloride and oxalic acid, both of which are soluble: CaC 2 4 t + 2HC1 1? CaCl 2 + H 2 C 2 4 . (1) CALCIUM AND ITS COMPOUNDS 395 The action of acids upon insoluble salts is so frequently mentioned in chemistry and is so important a factor in analytical operations that it demands separate discussion. This example is a typical one and may be used as an illustration. According to the rules already explained (p. 392), calcium oxalate (or any other salt) is precipitated when the numerical value of the product of the concentrations of the two requisite ions [Ca^] X [C2C>4 = ] exceeds the value of the ion-product for a saturated solution of calcium oxalate in pure water, that is, ex- ceeds 0.0 8 185 (p. 393). When, on the contrary, the product of the concentrations of the two ions falls below the limiting value, a condition which may arise from the removal in some way either of the Ca*" 1 " or of the C2C>4 = ions, the undissociated molecules will ionize, and the solid will dissolve to replace them until the ionic concentrations necessary for equilibrium with the molecules have been restored or until the whole of the solid present is consumed. Here the oxalate-ion from the calcium oxalate combines with the hydrogen-ion of the acid (usually an active one) which has been added, and forms molecular oxalic acid: + 2H+ fc* H 2 C 2 O 4 . (2) Hence, dissociation of the dissolved molecules of calcium oxalate proceeds, being no longer balanced by encounters and unions of the now depleted ions, and this dissociation in turn leads to solu- tion of other molecules from the precipitate. It will be seen that the removal of the ions in this fashion can result in considerable solution of the salt only when the acid pro- duced is a feebly ionized one. Here, to be specific, the concentra- tion of the C 2 O 4 = in the oxalic acid equilibrium (2) above must be less than that of the same ion hi a saturated calcium oxalate solution. Now oxalic acid does not belong to the least active class of acids, and its pure solution contains a considerable con- centration of C2O4 = . There is, however, a decisive factor in the situation which we have not yet taken into account. The hydro- 396 TEXT BOOK OF CHEMISTRY chloric acid which we used for dissolving the precipitate is a very highly ionized acid and gives an enormously greater concentration of hydrogen-ion than does oxalic acid. Hence the hydrogen-ion is in excess in equation (2), and the condition for equilibrium, rTT+12 vy [f^ O 1 - - - - . = K, will be satisfied by a correspondingly small [H 2 C 2 O 4 ] concentration of C 2 4 = . In this particular case, therefore, the [C 2 O 4 = ] of the oxalic acid is less than that given by the calcium oxalate. The whole change, therefore, depends for its accomplish- ment, not only on the mere presence of hydrogen-ion, but on the repression of the ionization of the oxalic acid by the great excess of hydrogen-ion furnished by the active acid that has been used. As a matter of fact, we find that a weak acid like acetic acid has scarcely any effect upon a precipitate of calcium oxalate. An acid stronger than oxalic acid must be employed. The whole scheme of the equilibria is as follows : CaC 2 4 (solid)<=CaC 2 O 4 (dslvd)*=Ca+++C 2 4 2HC1 = 2C1~~ -|- 2H + When excess of an acid sufficiently active to furnish a large con- centration of hydrogen-ion is employed, the last equilibrium is then driven forward and the others follow. With addition of a weak acid, only a slight displacement occurs, and the system comes to rest again when the molecular oxalic acid has reached a sufficient concentration. A generalization may be based on these considerations: an in- soluble salt of a given acid will in general interact and dissolve when treated with a solution containing another acid, provided that the latter acid is a much more highly ionized (more active) one than the former (see below). But even active acids frequently fail to bring salts of weak acids into solution, especially when the weak acid is itself present also. Here the cause lies in the fact that such salts are even less soluble than those of the calcium oxalate type, and give so low a con- CALCIUM AND ITS COMPOUNDS 397 centration of the negative ion that the utmost repression of the ionization of the corresponding acid does not give a lower value for the concentration of this ion than does the salt itself. Thus, we have seen that even hydrochloric acid (dilute) will not dissolve a number of sulphides. For example, in the case of cupric sulphide in a solution saturated with hydrogen sulphide, the S = factor in the solubility product [CU++] X [S = ] remains smaller than that in the scheme defining the hydrogen sulphide equilibrium [H + ] 2 X [S = ] even when the [S = ] factor in the latter is diminished in consequence of great addition of hydrogen-ion. In this case the first link in the chain of equilibria : CuS (solid) 5 CuS (dslvd) fc* Cu++ + S = ) 2HC1 } ' tends so decidedly backward that only the use of concentrated acid will increase the concentration of the H + to an extent sufficient to secure even a slight advance of the whole action. We must add, therefore, to the above rule : provided also that the salt is not one of extreme insolubility. This point will be illustrated more fully in connection with the description of individual sulphides (see under Cadmium). Illustrations of the application of these generalizations are countless. Carbonic acid is made from marble (p. 198), hydrogen sulphide from ferrous sulphide (p. 268), hydrogen peroxide from sodium peroxide (p. 303), and phosphoric acid from calcium phos- phate (p. 339). In each case the acid employed to decompose the salt is more active than the acid to be liberated. On the other hand, calcium oxalate is insoluble in acetic acid because this acid is weaker than is oxalic acid. We have thus only to examine the list of acids showing their degrees of ionization (p. 170) in order to be able to tell which salts, if insoluble in water, will be dis- solved by acids and, in general, what acids will be sufficiently active in each case for the purpose. In chemical analysis we dis- criminate between salts soluble in water, those soluble in acetic 398 TEXT BOOK OF CHEMISTRY acid (the insoluble carbonates and some sulphides, FeS and MnS, for example), those requiring active mineral acids for their sorti- tion (calcium oxalate and the more insoluble sulphides, for ex- ample), and those insoluble in all acids (barium sulphate and other insoluble salts of active acids). Precipitation of Insoluble Salts in Presence of Acids. The converse of solution, namely, precipitation, depends upon the same conditions : an insoluble salt which is dissolved by a given acid cannot be formed by precipitation in the presence of this acid. Thus calcium oxalate can be precipitated in presence of acetic acid, but not in presence of active mineral acids in ordinary concentrations. Cupric sulphide or barium sulphate can be precipitated in pres- ence of any acid, but ferrous sulphide and calcium carbonate only in the absence of acids. From this it does not follow that calcium oxalate, for example, cannot be precipitated if once an active acid has been added to the mixture. To secure precipitation, all that is necessary is to re- move the excess of hydrogen-ion which is repressing the ionization of the oxalic acid. This can be done by adding a base, which re- moves the H + , or even by adding sodium acetate. The acetate- ion C 2 H 3 O 2 ~ unites with the H + to form the little ionized acetic acid, in presence of which calcium oxalate can be precipitated. Basic Oxides. An oxide which combines (directly or in- directly) with water, giving a base, is called a basic oxide. Such an oxide (e.g. CaO) is always an oxide of a metallic element. All oxides of metallic elements are not basic, however (e.g. Mn 2 O7 and CrOa, pp. 498, 499, are acid anhydrides). The basic oxide is often used instead of the base, and gives the same compounds. For example, acids act upon calcium oxide., giving salts, just as they do with calcium hydroxide (p. 389) CaO + 2HC1 - CaCl 2 + H 2 O. Ca(OH) 2 + 2HC1 -> CaCl 2 + 2H 2 O. CALCIUM AND ITS COMPOUNDS 399 Acid anhydrides also combine directly with basic oxides to give the same salts which they form when acting upon the hydroxide. CaO + C0 2 -* CaC0 3 . Ca(OH) 2 + C0 2 -> CaC0 3 + H 2 0. Calcium Sulphate CaSOi. Various Forms. Calcium sul- phate is a very common mineral. It occurs, as anhydrite CaSO 4 , in the salt deposits. Gypsum CaSO 4 ,2H 2 O is found in masses, and also in single crystals (selenite, Fig. 42, p. 111). Alabaster is highly crystalline gypsum, tinted by impurities. Properties and Uses. Gypsum CaSO 4 ,2H 2 is the com- monest form, and is the one produced when calcium sulphate is precipitated. It is white and much softer than calcite. It is only slightly soluble in water (1 : 500 at 18). It is used in mak- ing plaster of Paris and is the chief component of blackboard " crayon." When gypsum is heated, the vapor pressure of the water it gives off soon exceeds that of the moisture in the atmosphere, and the compound begins to decompose: 2CaS0 4 ,2H 2 <= (CaSO 4 ) 2 ,H 2 + 3H 2 | . The hemi-hydrate which remains (plaster of Paris) gives a much lower pressure of water vapor and is more stable. Plaster of Paris (CaSO 4 ) 2 ,H 2 O. The hemi-hydrate is manu- factured in large quantities by heating gypsum in kilns. When moistened with water, it sets in about half an hour to a solid mass of gypsum. The temperature used in making it must not exceed 125, otherwise the hemi-hydrate is itself decomposed, the plaster is "dead burnt," and it no longer sets readily. The setting involves, simply, the reversal of the equation given above. Plaster of Paris swells somewhat, in setting, and so fills out completely every detail of a mould and applies itself closely to the outline of an object on which it is spread. It is used in mak- 400 TEXT BOOK OF CHEMISTRY ing casts, and in surgical bandages where movable parts are to be held rigidly in place. Stucco is made with sizing or glue instead of pure water. Casts are made smooth and non-porous (" ivory" surface) by a coating of paraffin which fills the pores. Excellent imita- tions of bronze or other castings are produced by rubbing with pulverized metals. Hard Water, Its Contents. All natural waters except rain water, which is "-soft," contain salts of calcium and magnesium in solution and are more or less "hard." These salts are dissolved by the water in its passage over and through the soil. Although limestone is very insoluble in pure water (0.013 g. per liter), yet it interacts with the carbonic acid contained in all natural waters, giving calcium bicarbonate which is about thirty times more soluble: C0 2 + H 2 * H 2 C0 3 + CaC0 3 Ca(C0 2 Ci5H 31 ) 2 i + NaJ30 4 . I Not until all the salts causing the hardness have been decom- posed, does the permanent solution of soap which is required for washing begin to be formed. This waste is often very great and expensive. Thus, calculating from the equation, with 1 (U.S.) jof hardness, 100 gallons (U.S.) of water should use up 0.075 i pound of soap (1 Brit, and 100 gal. Brit., 0.075 lb.). In point 402 TEXT BOOK OF CHEMISTRY of fact, however, twice these equivalent amounts of soap are wasted, because the colloidal calcium salts carry down with them more than an equal amount of undecomposed soap. Hence measurement shows that, with 1 (U.S. or Brit.) of hardness, 100 gallons (U.S. or Brit.) of water actually destroy 0.17 pound of soap. Thus, with 35 no less than 6 pounds of soap are wasted per 100 gallons, before the part of the soap that is to do the work begins to dissolve. Treatment of Hard Water. The temporary hardness can be removed by boiling the water, or using some preheating arrange- ment in connection with the boiler (stationary engines only). Temporary hardness is commonly removed, on a large scale, by adding slaked lime (made into milk of lime) in exactly the quantity shown by an analysis of the water to be required, and stirring for a considerable time: Ca(HCO 3 ) 2 + Ca(OH) 2 -* 2CaC0 3 I + 2H 2 0. (1) The bicarbonate is neutralized and all the lime precipitated. The latter is removed by filtration. Permanent hardness is not affected by slaked lime, but is pre- cipitated by adding sodium carbonate in the necessary proportion: CaS0 4 + Na 2 C0 3 -* CaC0 3 I + Na 2 S0 4 . (2) When both kinds of hardness are present, crude caustic soda (sodium hydroxide) may be employed. It neutralizes the bi-; carbonate, precipitating CaC0 3 : Ca(HCO 3 ) 2 + 2NaOH -> CaC0 3 | + Na 2 C0 3 + 2H 2 (3) and giving sodium carbonate. The latter then acts as in equation (2). Instead of this, the treatments indicated in equations (1) and] (2) may be applied in combination (Porter-Clark process).* * So far as the hardness is due to magnesium bicarbonate, a double propor- tion of lime must be added to precipitate the magnesium as hydroxide (sol'ty 0.01 g. per 1.), because the carbonate is too soluble (1 g. per 1.). CALCIUM AND ITS COMPOUNDS 403 In the new permutite process the water is simply filtered through an artificial sodium silico-aluminate (permutite NaP) which is supplied in the form of a coarse sand. The calcium, etc., in the water is exchanged for sodium, which does no harm : Ca(HC0 3 ) 2 + 2NaP -> 2NaHC0 3 + CaP 2 . After twelve hours' use, the permutite is covered with 10 per cent salt solution and allowed to remain for the other twelve hours of the day, when it is ready for employment once more : 2NaCl + CaP 2 - CaCl 2 + 2NaP. Only salt, which is inexpensive, is consumed, and calcium chloride solution is thrown away. Permutite removes magnesium, iron, manganese, and other elements in the same way. The life of a charge is said to be over twenty years. Hard Water in the Laundry. As we have seen (p. 377), soap will soften water, but the calcium and magnesium salts of the soap acids, which are precipitated, are sticky, and soil the goods being washed. Other substances that soften water, not only give non-adhesive precipitates, but are also much cheaper, and an attempt is generally made to utilize them. The use of slaked lime is impracticable on a small scale. Washing soda Na 2 CO 3 ,10H 2 O is added to precipitate both kinds of hardness: Ca(HC0 3 ) 2 + Na 2 C0 3 -> CaCO 3 I + 2NaHCO 3 CaSO 4 + Na2CO 3 -> CaCO 3 | The small amounts of salts of sodium which remain in the water have no action on soap. Household Ammonia NH 4 OH acts like sodium hydroxide (p. 404) : Ca(HCO 3 ) 2 + 2NH 4 OH - CaCO 3 j + (NH 4 ) 2 C0 3 + 2H 2 CaSO 4 + (NH 4 ) 2 CO 3 - CaCO 3 j + (NH 4 ) 2 SO 4 except that it will not precipitate magnesium-ion. 404 TEXT BOOK OF CHEMISTRY Borax Na 2 B 4 O7,10H 2 is hydrolyzed (p. 349) and the sodium hydroxide contained in its solution acts as already (p. 349) de- scribed. The supposed bleaching or whitening action of borax or soda is a myth; these salts prevent straining by the iron in the water. They simply precipitate the iron (present as Fe(HC0 3 )2), which almost all waters contain, as FeC0 3 before the goods are put in. This precipitate is easily washed out in rinsing. The palmitate, etc., of iron, however, which the soap itself would throw down, is sticky and adheres to the cloth. The air. subsequently oxidizes it and gives hydrated ferric oxide (rust), which is brownish-red. It is evident that, properly to achieve their purpose, the soda and borax must be added, must be completely dissolved, and must be allowed to produce the precipitation of FeCO 3 , CaC0 3 , etc., all before the soap, or the goods, are introduced. If the soap is dissolved before or with the soda, it will take part in the pre- cipitation, and give sticky particles containing the iron and cal- cium salts of the soap acids. The soda, borax, and ammonia do not themselves remove dirt that is done by the dissolved soap (p. 379). With the help of rubbing, however, they do emulsify and remove animal or vege- table oil and grease, but not mineral oil (pp. 309 and 310), when these happen to be on the goods. But soap alone will do this also, and remove mineral oil as well. Washing powders are, or ought to be, mainly sodium carbonate, mixed with more or less pulverized soap. Calcium Chloride, CaCl 2 . Chloride of calcium is obtained as a by-product in the Solvay process (p. 352) and in other indus- tries. It crystallizes from water as the white hexahydrate, CaCl 2 ,6H 2 O, and is very soluble. It has few applications. The porous, granular variety, used for drying gases, is made by driving most of the water out of the hexahydrate by heat. The granular form is used in large amounts for sprinkling on dusty roads. CALCIUM AND ITS COMPOUNDS 405 The salt, being deliquescent (p. 404), attracts water from the air and moistens the dust with calcium chloride solution. The satu- rated solution does not freeze until 48 is reached, so that chilled calcium chloride brine is used in refrigerating appliances (p. 261). Calcium Cyanamide CaCN 2 . Calcium carbide, when strongly heated, absorbs nitrogen, giving a mixture of calcium cyanamide and carbon (nitre-lime) : CaC 2 + N 2 -* CaCN 2 + C. When treated with hot water, the product is decomposed into calcium carbonate and ammonia: CaCN 2 + 3H 2 -> CaC0 3 + 2NH 3 . Calcium cyanamide is now manufactured for use as a fertilizer. In the moist, cold soil the action is probably not so simple as that shown in the above equation, but it finally pro- vides combined nitrogen in a form that can be absorbed by plants. At Odda (Norway) the carbide is pulverized and placed in a cylindrical furnace (Fig. 96), holding 300 to 450 kg. The heat (800 to 1000) is fur- nished by the passage of a current of electricity through a thin carbon rod, which passes through the axis. The tube surrounding the rod and the other ^r? I c partitions are of cardboard, which burns up and leaves openings for the circulation of the nitrogen. The latter is made by the fractionation of liquid air and is introduced under pressure. In thirty-five hours nitrogen ceases to be absorbed, and the product is pulverized when cold. Nitro-lime, when fused with sodium carbonate, gives sodium cyanide NaNC, used in the extraction of gold: CaCN 2 + C + Na^COs -> CaCO 3 + 2NaNC. 406 TEXT BOOK OF CHEMISTRY Other Compounds of Calcium. Calcium fluoride CaF 2 occurs as a mineral (fluorite). It is our source of hydrofluoric acid (p. 332), and is used in metallurgy to lower the melting-point of slags. The phosphates (p. 339) and bisulphite of calcium, and bleaching powder CaCl(OCl) (p. 307) have already been discussed. Exercises. 1. Why can not lime-water be kept in an open bottle? Lfev> +&OK- M = e*e* a *w 2. Why does whitewash become so firmly attached to the wall? 3. When the water drips from the tap into a porcelain basin, a brown stain is gradually produced. Explain. Make equations for the action, assuming the product to be Fe20 3 . 4. Make equations for the action of sodium palmitate: (a) upon calcium bicarbonate; (b) upon magnesium sulphate. 5. In softening water: (a) what would be the objection to using an excess of milk of lime; (b) why is prolonged stirring required (p. 403); (c) why must the precipitate be removed by filtration ? 6; (a) Why do small amounts of sodium salts have no action on soap? (b) What effect would large amounts have? 7. Explain why wood ashes are sometimes used to soften water, and how they act. 8. Make equations for the action of chlorine upon quicklime. 9. Why does fluorite lower the melting-point of a slag? IONIC EQUILIBRIUM, CONSIDERED QUANTITATIVELY In view of the predominance of ionic actions in the chemistry ot the metals, and of the determinative effect of ionic equilibria on many actions, it is essential that we should be prepared in future for a more exact consideration of these phenomena than we have hitherto attempted. The whole basis for this exact consideration has already been supplied, and only more specific application of the principles is demanded. The basis referred to, which should now IONIC EQUILIBRIUM, CONSIDERED QUANTITATIVELY 407 be re-read as a preliminary to what follows, is contained in, (1) the discussion of chemical equilibrium in general (pp. 228-241), (2) the application of the same principles to ionic equilibrium (p. 391), and (3) the illustration of this application in the case of cupric bromide (p. 407). Excess of One Ion. In the case of cupric bromide, we showed that increasing the concentration of the bromide ions displaced the equilibrium by favoring the union of the ions to form molecular cupric bromide: 2Br~ -f Cu++ -> CuBr 2 . This we speak of as a repression of the ionization of the cupric bromide. Now, if the sub- stance is a slightly ionized one, like a weak acid or a weak base, the repression of the ionization through the formation of molecules in this way may remove so many of that one of the ions which is not present in excess (corresponding to the Cu++ in the foregoing illus- tration), that the mixture will no longer respond to tests for the ion so removed. This is an interesting and very common case. The behavior of acetic acid, a weak, slightly ionized acid, will serve as an illustration. In normal solution (60 g. in 1 1.) acetic acid is only 0.004 ionized (p. 170), so that, in the equation for the equilibrium, (0.996) HC 2 H 3 2 1=> H+ (0.004) + C 2 H 3 2 ~ (0.004), the relative proportions are as shown by the numbers in paren- theses. If the whole of the acid (60 g.) were ionized, there would be 1 g. of hydrogen-ion per liter. Yet, even in the much smaller concentration actuaUy present (0.004 g. per liter), the acid taste of the H+ and its effect upon indicators can be distinctly recognized. If, now, solid sodium acetate is dissolved in the solution, the liquid no longer gives an acid reaction with one of the less delicate indica- tors, like methyl orange. The explanation is simple. Sodium acetate is highly ionized. It gives, therefore, a large concen- tration of acetate-ion to a liquid formerly containing very little. [This causes a greatly increased union of the H+ ions and C 2 H 3 2 ~ 408 TEXT BOOK OF CHEMISTRY ions to occur, and the former, being already very few in number, disappear almost entirely. Hence the solution becomes, to all intents and purposes, neutral. There is no less acetic add present than before, but the concentration of hydrogen-ion is very much smaller. Formulation and Quantitative Treatment of the Case of Excess of One Ion. If the semi-mathematical mode of formu- lating an equilibrium (p. 394), as applied to the case of an ionogen (p. 395), be employed here, the foregoing general statements may be made more precise and the conclusions clearer. If [H + ] and [C 2 H 3 2 ~] represent the molecular concentrations of hydrogen-ion and acetate-ion, respectively, and [HC 2 H 3 2 ] that of the acetic molecules at equilibrium, then: [H+] X [C 2 H 3 2 -] _ [HC 2 H 3 2 ] The value of K is constant, whether the strength of the solution acetic acid is great or small, and even when another substance with a common ion is present. In the latter case, [C 2 H 3 O 2 ~] and [H + ] stand for the whole concentrations of each of these ionic substances from both sources. Now, in normal acetic acid [H+] = 0.004, [C 2 H 3 O 2 ~] = 0.004 (for the number of each kind of ions is the same), and [H^HsC^] = 0.996, practically 1. Substituting in the formula: 0.004 X0.004 = g( = 0()4l6) 1 When, however, sodium acetate is dissolved in the liquid until the solution is normal in respect to this substance also, the following additional equilibrium has to be considered : (0.47) NaC 2 H 3 O 2 <= Na+ (0.53) + C 2 H 3 2 ~ (0.53). The concentration of acetate-ion from this source is 0.53, so that, in the mixture of acid and salt, the concentration of acetate-ion IONIC EQUILIBRIUM, CONSIDERED QUANTITATIVELY 409 [C 2 H 3 O 2 ~] will be 0.53 -f 0.004 = 0.534, or nearly 134 times larger than in the acid alone. Hence, in order that the product [H+] X [C 2 H 3 O 2 ~] may recover, as it must, a value much nearer to the old one, [H + ] must be diminished to something like yi? of its former magnitude. That is, [H + ] will become equal to about 0.00003, 000008 X 0.534 ' the rest of the hydrogen-ion uniting with a corresponding amount of the acetate-ion to form molecular acetic acid. The effect of adding this amount of sodium acetate therefore is, as we have seen, to reduce the concentration of the hydrogen-ion below the amount which can be detected by use of an indicator like methyl- orange. This effect is of course reciprocal, and the ionization of the sodium acetate will be reduced also. But the acetate-ion furnished by the acetic acid is relatively so small in amount (0.00003 against 0.53) that the effect it produces on the ionization of the salt is imperceptible. It will be noted that the acetate-ion and hydrogen-ion disappear in equivalent quantities, for they unite. There is, however, so much of the former that the loss it sustains goes unremarked, while there is so little of the latter that almost none of it remains. When substances of more nearly equal degrees of ionization are used, both effects are equally inconspicuous. Thus, sodium chloride and hydrogen chloride hi normal solutions yield approximately equal concentrations of chloride-ion (0.784 and 0.676). Hence, if one mole of sodium chloride were to be dissolved in the portion of water already containing one mole of hydrogen chloride, the concentra- tion of the chloride-ion, at a very rough estimate, would be nearly double. If this doubling of the concentration of chloride-ion almost halved that of the hydrogen-ion (0.784), in order that the expression [CF] X [H + ] -f- [HC1] might remain constant, the concentration of the hydrogen-ion would still be about 0.400 and 410 TEXT BOOK OF CHEMISTRY therefore 100 times as great as in molar acetic acid. It is thus altogether impossible to reduce the concentration of the hydrogen- ion given by an active acid like hydrochloric acid below the limit at which indicators are affected, for there is no way of introducing the enormous concentration of the other ion which the theory demands. With more crude means of observation than indicators afford, effects like this last may sometimes be rendered visible. This was the case with cupric bromide solution, to which potassium bromide was added (p. 407). The blue of the cupric-ion disappeared from view, while much cupric-ion was still present, because the brown color of the molecular cupric bromide covered it up completely. Special Case of Saturated Solutions. The commonest as well as the most interesting application of the conceptions de- veloped above is met with in connection with saturated solutions, especially those of relatively insoluble substances. The situation in a system consisting of the saturated solution and excess of the solute has been discussed already (read p. 116). In the case of potassium chlorate, for example, we have the follow- ing scheme of equilibria: KC10 3 (solid) 2ZnO + 2SO 2 . The ore is fed in at the top of a huge, box-like furnace (Fig. 98, diagrammatic) through which rush the flames and heated from fuel gas burning with an excess of air. The ore is turned over, and gradually displaced forward by moving rakes until, at the end, it drops to the next level. Here it is raked in the opposite direction, until it falls to the third level. The ore collects at the bottom fully oxidized, while the sulphur dioxide in the gases is made into sulphuric acid. The oxide from either ore is then reduced by heating with powdered coal: C->Zn 418 TEXT BOOK OF CHEMISTRY This treatment of zinc ores should be carefully considered. Since ores of most metals consist of the carbonate, sulphide, or oxide of the metal, these steps are common to most metallurgical processes. In the case of other metals, only the forms of the furnaces and other details vary. In the case of zinc, because it is a volatile metal, the heating of the mixture of oxide and coal is conducted in retorts, from which the metal issues as vapor. The mixture is placed in fire-clay cylinders (4 to 5 ft. long), which are arranged in several tiers in an oblong, gas-heated furnace (Fig. 99). A fire-clay receiver is FIG. 99. luted on with clay. The carbon monoxide burns with a blue flame at the nozzle of each receiver, while the zinc condenses to liquid within it. From time to time the liquid metal is raked out into a traveling iron pot, from which it is poured into moulds. Properties and Uses. The metal is silvery and crystalline. At 120 to 150 it can be rolled into sheets between hot rollers, at 200 to 300 it becomes brittle, at 419 it melts, and at 950 it boils. The density of the vapor shows it to be monatomic. Zinc vapor burns with a bluish flame, giving ZnO. In air the metal does not rust, being protected by a non-porous coating of a basic carbonate which adheres closely to the surface. Zinc displaces hydrogen from dilute acids (p. 54). ZINC zn 419 Sheet zinc is used for gutters and cornices, although for the latter purpose it is being displaced by terra cotta. Iron is coated (galvanized) with zinc by thorough cleaning with dilute sulphuric acid or the sand blast, and dipping in melted zinc. Iron netting, corrugated iron for sheds and roofing, and iron gutters, tanks, and pipes are coated, either in this way, or galvanically (see copper). Sherardized iron is made by covering the article with zinc dust and baking it. The zinc protects the iron, primarily because it ex- cludes the air from the surface, and secondarily because, even when the coating is broken, the zinc, being the more active metal of the two (p. 54), is rusted first. Zinc is used also in extracting silver from crude lead (p. 488), as the active metal (anode) in batteries, and in several alloys (e.g. Babbitt's metal, p. 342, brass, and German silver). In the laboratory granulated zinc, made by pouring the melted metal in a thin stream into water, and zinc dust (impure, contains ZnO), are the forms commonly employed. Zinc Oxide ZnO and Hydroxide Zn(OH) 2 . The oxide is made by burning zinc vapor in air. It is yellow while hot, and white when cold. Mixed with oil, it is used as a paint (Chinese white). It has less covering power than has white lead paint (4 coats of the former equal 3 of the latter), but it does not darken from exposure to hydrogen sulphide in the air (ZnS is white, PbS black). Zinc hydroxide Zn(OH) 2 is formed by precipitation. Both the oxide and hydroxide are basic, and give salts with acids. But they are also somewhat acidic, and dissolve in excess of sodium hydroxide, giving sodium zincate NaHZnC>2. When heated with a salt of cobalt, they give the green cobalt zincate CoZn02 (test). Zinc Chloride ZnCl 2 . The chloride is* formed in the action of zinc or zinc oxide on hydrochloric acid. It is a white, deliques- cent solid. Its aqueous solution gelatinizes cellulose and dissolves it (p. 227), and thus is used in parchmentizing paper and in im- 420 TEXT BOOK OF CHEMISTRY pregnating wood to prevent decay. The aqueous solution is acid (hydrolysis) and is used for cleaning the surface of metals before soldering. Solder "runs" on a hot brass, copper, or lead surface, provided the latter is clean, and adheres perfectly when cold. But it does not dissolve oxides, or melt them, and therefore cannot even reach the surface of the metal, much less adhere to it, if the slightest tarnish is present. Other Compounds. Zinc sulphate, ZnS0 4 , 7H 2 0, is made by the action of sulphuric acid on zinc or zinc oxide. It is used in preserving hides and as a mordant in cotton printing (see dyeing, p. 430). Zinc sulphide ZnS is precipitated (white) when ammo- nium sulphide (NH 4 ) 2 S solution is added to a salt of zinc: ZnS0 4 + (NH 4 ) 2 S<=ZnS j + (NH 4 ) 2 SO 4 (dslvd.). All double decompositions are reversible, but this one is practically complete. This is because ammonium sulphide, being a salt, is highly ionized and gives much sulphide-ion S = . When hydrogen sulphide is employed instead ZnS0 4 + H 2 S + ZnS + H 2 SO 4 only a part of the zinc is precipitated. The action comes to rest because the hydrogen-ion H + of the sulphuric acid competes with the zinc-ion Zn++ for the S = , to form molecular H 2 S, and when a certain concentration of sulphuric acid (and therefore of H+) has been reached, no more ZnS can be formed. A mixture of zinc sulphide and barium sulphate BaS0 4 , pre- pared in a special way, is called lithopone. Used as a white pigment, it has greater covering power than has white lead and is, besides, non-poisonous. All the soluble compounds of zinc are somewhat poisonous. Exercises. 1. (a) How do the electrolytic methods of making calcium and magnesium differ ? (6) Why not electroly ze an aque ZINC zn 421 ous solution of magnesium chloride in making magnesium? (c) Why use both potassium and sodium chlorides in making magne- sium? (d) Why is magnesium, but not potassium or sodium, liberated? 2. Why are magnesium and zinc not found free in nature? 3. Why does magnesium rust completely (in time), while zinc does not? 4. Make equations for: (a) the action of magnesium on hydro- chloric acid; (b) the burning of magnesium in air; (c) the heating of magnesium carbonate; (d) the precipitation of magnesium hydroxide from the sulphate; (e) the hydrolysis of magnesium chloride. 5. Why is salt containing magnesium chloride, after mixing with sodium bicarbonate, no longer deliquescent? 6. What is the density of zinc vapor (air = 1) ? 7. Make equations for: (a) the action of hydrogen sulphide on zinc oxide; (b) the precipitation of zinc hydroxide; (c) the action of sodium hydroxide on zinc hydroxide. CHAPTER XXXVI ALUMINIUM THE family of trivalent elements to which the metal aluminium belongs includes the non-metal boron (p. 348), and several rare metals. Occurrence. Aluminium, although it does not occur free, is the third element in order of quantity. Its silicates, such as clay (kaolin) HAlSi0 4 , mica KAlSi0 4 , and felspar KAlSi 3 8 , are amongst the most plentiful minerals. The oxide A1 2 3 occurs as corundum, sapphire, ruby, and emery (impure form). Bauxite is a valuable hydrated oxide. Garnets Ca 3 Al 2 (Si04) 3 are mined and pulverized to make "sand" paper. Cryolite 3NaF,AlF 3 (Greek, ice-stone) is imported from Greenland. Manufacture. The making of aluminium on a large scale originated in C. M. Hall's discovery (1886) that the oxide could be electrolyzed in solution in molten cryo- lite. Iron boxes (Fig. 100), about 5 by 3 feet and 6 inches deep, are lined _ _-. with a compressed mixture of coke and J I I 1 1 B I tar which is afterwards baked. The lining forms the cathode, while the . jj oxygen is liberated at the anodes ^^^^| a series of rods of carbon about 3 inches in diameter which are attached The cryolite is melted (1000) by the arcs The latter are then raised some- FIG. 100. to copper rods, struck by the carbon rods, what, the aluminium oxide is added, and some coal (which floats) is thrown in to cover the surface and obscure the blinding glow. 422 ALUMINIUM 423 From time to time more of the oxide is added and the melted alu- minium (m.-p. 650 to 700) is tapped off. The oxide must be made from carefully purified bauxite, as the metal itself can not be purified commercially. In 1866 it cost $250 to 750 per kilogram and now sells at about 50 cents (2/-). Physical Properties. The metal has a lower specific gravity (2.6) than any other metal that could be put to the same uses (sp. gr. iron 7.8, copper 8.8). It has malleability and the foil is taking the place of tinfoil to some extent for wrapping foods. It has considerable tensile strength, and is a good conductor of electricity. It is lacking in sufficient hardness and, when pure, can not be filed or turned because it adheres to the tools. Chemical Properties. The metal displaces hydrogen rapidly from hydrochloric acid, it acts very slowly on dilute sulphuric acid, and hardly at all on cold nitric acid. In the air it acquires only a slight film of closely adhering oxide. This film prevents it from acting upon water (hot or cold). When the surface is cleaned and amalgamated with mercury, by dipping in mercuric chloride solution, however, this metal acts as a contact agent, and hydrogen is rapidly displaced: 2A1 + 6H 2 O-+3H 2 t + 2A1(OH) 3 I . The foil, when heated, burns brilliantly in the air. When amal- gamated (as above), it rusts in moist air with astonishing rapidity. Uses. The largest quantity is consumed by steel-makers. When added in small amount (less than 1 : 1000) to molten steel, it combines with the gases, and gives sound ingots free from blow holes. Next to this comes its use for long distance transmission of electricity. A cable of the requisite capacity is larger than one of copper for the same current, but is lighter and puts less strain on the supports. Cooking vessels of aluminium are not corroded 424 TEXT BOOK OF CHEMISTRY and are largely used. Cameras and opera glasses are made of it. Pulverized aluminium, mixed with oil, gives a paint which protects iron admirably. Aluminium bronze (copper, with 5 to 12 per cent aluminium) has a brilliant golden yellow color and is stable in air and easily worked. Magnalium (containing 1 to 2 per cent of magnesium) can be filed, turned, or polished like a mirror, and is better for many purposes than the pure metal. On account of its great chemical activity (p. 54), aluminium dis- places many other metals (e.g. iron, manganese, and chromium) from combination (Goldschmidt, 1898). Thus powdered alumin- ium and oxide of iron when mixed (thermite) in a crucible, and started by a burning magnesium ribbon, interact with great violence : 2A1 + Fe 2 3 - A1 2 3 + 2Fe. A temperature of 3000 to 3500 is reached, the molten iron (m.-p. 1700) collects at the bottom, and the molten aluminium oxide (m.-p. 2050) floats to the top. Steel rails are welded together in situ, and large objects of steel like broken propeller shafts are mended, by enclosing a mass of thermite round the joint and firing it. Aluminium Hydroxide A1(OH) 3 . The hydroxide is pre- cipitated when ammonium hydroxide, or other alkaline hydroxide, is added to a solution of a salt of aluminium : A1 2 (S0 4 ) 3 + 6NH 4 OH -> 2A1(OH) 3 j + 3(NH 4 ) 2 S0 4 (dslvd.). It tends to remain in colloidal suspension, and forms a white gelatinous precipitate. The mineral bauxite contains a smaller proportion of the elements of water. Aluminium hydroxide, when heated, leaves the oxide: 2A1(OH) 3 -+ A1 2 3 + 3H 2 O | . ALUMINIUM 425 It is both weakly basic and feebly acidic in chemical properties. In acids it dissolves forming salts of aluminium, such as the chloride or acetate: A1(OH) 8 + SHCl^AlCla + 3H 2 0. A1(OH) 3 + 3HC0 2 CH 3 A1(C0 2 CH 3 ) 3 + 3H 2 0. These actions are not complete and the salts are hydrolyzed in water (read the equations backwards) giving acid solutions (p. 356). The hydroxide acts also as a feeble acid towards strong bases (see p. 425). Aluminium hydroxide, precipitated from the sulphate, is used in sizing paper (to fill the capillaries and pores), in purifying water (see p. 427), and as a mordant (see p. 430) in dyeing. Delicate fabrics (cloth) are rendered waterproof by saturating them with aluminium acetate solution and boiling to promote the hydrolysis. The aluminium hydroxide is precipitated in the capillaries of the cotton or linen, rendering them non-absorbent Alumina tes. Aluminium hydroxide dissolves in sodium hydroxide solution, to form sodium aluminate Na 3 AlO 3 : 3NaOH + H 3 A1O 3 -> Na 3 AlO 3 + 3H 2 and when fused with calcium oxide (or carbonate) it gives calcium aluminate : Skeleton: CaO + H 3 A10 8 -> Ca 3 (A10 3 ) 2 + H 2 O. Balanced: 3CaO + 2H 3 A10 3 -> Ca 3 (A10 3 ) 2 + 3H 2 O. The aluminates are hydrolyzed by water, and their solutions have an alkaline reaction. Compounds of aluminium, when heated with a salt of cobalt, leave a blue cobalt aluminate (Co) 3 (A10 3 ) 2 (test). Aluminium Oxide (Alumina) A1 2 O 3 . Corundum, and the impure variety emery, are next to the diamond in the scale of hardness, and are used as abrasives, Euby and sapphire are also 426 TEXT BOOK OF CHEMISTRY crystallized aluminium oxide, containing traces of impurities (chromium in the one case and iron and titanium in the other) to which they owe their colors. By ingenious methods of fusing the oxide, " synthetic" sapphires to the extent of six million carats and rubies to the extent of ten million carats are now made annu- ally. The artificial gems are chemically identical with the natural ones, and can be distinguished only by the fact that they are free from microscopic bubbles and other defects. Alundum, an arti- ficial abrasive, and refractory material for crucibles and muffles, is made by barely melting the oxide in the electric furnace. Aluminium oxide is both basic and acidic. Aluminium, Sulphate Al 2 (SOi) 3 ,18H 2 O. The sulphate is manufactured by the action of sulphuric acid on bauxite. It crystallizes in leaflets, which usually have a faint yellow tinge due to the presence of iron (Fe(OH) 3 ) derived from the mineral. The salt is used in fireproofmg cloth, since, when heated, it melts in its water of hydration. It is used as a source for precipitated alu- minium hydroxide in paper-making, water purification, and dyeing. Alums. When aluminium sulphate and potassium sulphate are dissolved together in molecular proportions, the solution deposits transparent octahedral (Fig. 39, p. Ill) crystals of potash-alum K 2 S0 4 , A1 2 (S0 4 ) 3 , 24H 2 0. This salt is more easily freed from impurities (e.g. compounds of iron) by recrystallization than is aluminium sulphate, and is therefore used instead of the latter in medicine, in dyeing (delicate shades), and to replace the cream of tartar (p. 354) in making baking powder. Sodium-alum, Na 2 S04,Al 2 (S04) 3 , 24H 2 O, ammonium-alum (NH 4 ) 2 S0 4 , A1 2 (S0 4 ) 3 , 24H 2 O and chrome-alum K 2 SO 4 , Cr 2 (S0 4 ) 3 , 24H 2 O are made in the same way, and crystallize in the same form. The first two are used as sources of aluminium hydroxide, and the last in the " fixing bath" to harden the gelatine on photographic films and plates. ALUMINIUM 427 Water Purification Coagulation Process. The sus- pended matter in water to be used for a domestic supply can be coagulated into larger particles by introducing a small amount of the gelatinous precipitate of aluminium hydroxide. These larger particles, which enclose also the greater part of the bacteria, settle rapidly and the process therefore permits the use of relatively small settling ponds. Aluminium sulphate, made from crude bauxite, and lime are added to the water: 3Ca(OH) 2 + A1 2 (S0 4 ) 3 -> 3CaS0 4 + 2A1(OH),. If the water has much temporary hardness, lime is not required: 3Ca(HC0 3 ) 2 + A1 2 (S0 4 ) 3 -> 3CaS0 4 + 2A1(HCO 3 ) 3 . The carbonate of aluminium, being a salt of both a very weak acid and a very weak base, if formed, would be instantly hydrolyzed: A1(HC0 3 ) 3 + 3H 2 -> Al(OH), | + 3H 2 CO 3 so that aluminium hydroxide is precipitated. The few remaining bacteria are destroyed by the addition of bleaching powder (p. 308). Crude ferrous sulphate FeS0 4 (copperas), being in many places cheaper (see p. 464) than aluminium sulphate, is often used instead of the latter. The lime precipitates ferrous hydroxide Fe(OH) 2 . This is quickly oxidized to the red ferric hydroxide Fe(OH) 3 , which coagulates the suspended matter. Clay and Pottery. Pure clay (p. 453) or kaolin is white. It is hydrogen-potassium silicate HKSiO 4 . Common clay contains impurities such as sand (silica), limestone, and compounds of iron. Both kinds are plastic when wet and can be moulded. When heated strongly the material shrinks (so that the products are porous) and becomes hard. Bricks, and tiling for roofs and drains, are made of common clay and, when red, owe their color to oxide of iron (Fe 2 3 ) . The firing is done with fuel gas in ovens or kilns 428 TEXT BOOK OF CHEMISTRY of brickwork. To glaze drain pipes and some bricks, salt is thrown into the kiln. The vapor of the salt produces a more fusible sodium-aluminium silicate, which fills the surface pores. Clay for fire-brick (infusible) must contain free silica, but no lime. China and porcelain are white, translucent and non-porous. They are made of pure clay to which a little of the more fusible felspar is added. After firing, the articles are dipped in water, in which the materials for the " glaze," namely finely, ground felspar and silica, are suspended. Having thus acquired a thin coating of these substances, they are fired again at a higher tem- perature and for a longer time. Colored decoration is done with materials which melt (third firing) to colored enamels (p. 347). Eggshell porcelain contains no clay and is practically a trans- lucent glass. Ultramarine. This material is made by heating together kaolin, sodium carbonate, sulphur, and charcoal, pulverizing the green mass, and heating it again with more sulphur. The product is used as laundry blueing, and in making blue-tinted paper. It is added also to correct the yellow shade of linen, starch, sugar (p. 222), and paper stock. Cement. Cement is made by heating limestone CaCOs, clay HAlSi04, and sand Si02, or a natural rock containing all three in the right proportions. Such a rock, made into cement by vol- canic heat, was quarried by the Romans near Naples and else- where, and its capacity for hardening even under water was utilized by them. The mixture, or pulverized natural rock, is moistened and fed slowly in at the upper end of an inclined (6) revolving cylinder (20 to 45 by 2 meters). The motion continu- ally turns over the thin layer, and exposes every particle to the heat of the air-blast, charged with pulverized coal, burning in the interior. The product slides out in a continuous stream at the lower end, and is pulverized by steel balls in a ball mill. ALUMINIUM 429 Cement is held to be a mixture of calcium silicate and calcium aluminate. The former is simply a filler. The latter is hydrolyzed by the water: Ca 3 (A10 3 ) 2 + 6H 2 - 3Ca(OH) 2 + 2H 3 A10 3 . The calcium hydroxide slowly crystallizes, connecting the particles of the calcium silicate. The aluminium hydroxide fills the inter- stices and renders the whole compact and impervious. Concrete is a mixture of cement with sand and crushed stone or gravel, all made into a paste with water. It sets to a solid mass, suitable for walks, and for the foundations, walls, and floors of buildings. Since no carbon dioxide from the air is required in the hardening process, it sets equally well under water (hence hydraulic cement), and is employed in constructing dams, levees, and the foundations of bridges. Reinforced concrete contains twisted rods or iron, embedded in the mass, and is much used in building construction. The quantity of cement made in the United States has risen in ten years from 26 million barrels (of 380 pounds) to 80 million barrels (1911). Blast-furnace slag, when pulverized and heated with limestone, has been found to yield an excellent quality of cement, and a valuable use has thus been found for what was formerly an annoy- ing encumbrance. Dyeing. Cotton and wool are tubular in structure (Fig. 2, p. 2), while silk is a transparent, solid filament. After the dyeing, the color must be, as far as possible, fast to rubbing, light, and washing. When cotton (cellulose, a carbohydrate) is dipped in a solution of most dyes, the latter can be washed out very easily. Wool and silk (proteins, see p. 432), on the other hand, are permanently dyed by this treatment. Most dyes are thus "direct" dyes on wool or silk very few are so on cotton. Insoluble Dyes. If the dye is insoluble, and can be produced by precipitation after the solution has filled the walls and capillary 430 TEXT BOOK OF CHEMISTRY of each fiber of the goods, then it cannot afterwards be removed by water. Indigo is of this kind. The cloth is saturated with a solution of indigo white Ci 6 Hi 2 N 2 02, and the oxygen of the air subsequently oxidizes this and deposits the insoluble indigo blue Ci 6 HioN 2 02. Indanthrene blue is applied in the same way as indigo, and is even less affected by light. Mordant Dyes. Colloidal and gelatinous substances like alu- minium hydroxide coagulate and precipitate many dyes, which are themselves colloids. Such a substance is called a mordant. The colored precipitate is called a lake, and is insoluble. By first saturating the cloth with hot aluminium acetate solution, or by using first alum and then ammonium hydroxide, the alumin- ium hydroxide can be precipitated within the fibers of the goods. When the goods are then dyed, the coloring matter is taken up by the mordant, with which it forms an insoluble lake, within the fibers. Aluminium hydroxide forms insoluble lakes with acid dyes, such as Alizarin (madder). Chromium hydroxide, iron hydroxide, tin hydroxide, and other gelatinous, insoluble hydrox- ides are used for the same purpose. For basic dyes, like Malachite green and Methyl violet, an insoluble salt of tannic acid (such as antimonyl tannate) is a suitable mordant. Direct Dyes. Most dyes, like Poncean 6R (p. 2), are direct dyes on silk and wool only. Congo-red dyes cotton also directly. This dye, which is a sodium salt of an organic acid, is a colloid, and is salted out within the goods by adding sodium sulphate to coagulate it (see soap, p. 378), and assist it in forming an insoluble material with the cotton. Exercises. 1. Make equations for the following actions: (a) aluminium on hydrochloric acid; (b) aluminium on mercuric chloride HgCl 2 ; (c) displacement of manganese from manganese dioxide by aluminium. 2. How must aluminium hydroxide be ionized so as to behave both as an acid and a base ? ALUMINIUM 431 3. Make the equation for heating calcium carbonate with: (a) aluminium hydroxide; (b) aluminium oxide. 4. Explain why the reaction of solutions of aluminates is alka- line (p. 425) and that of alums is acid. 5. Make the equation: (a) for the action of sulphuric acid on bauxite, assuming the formula of the latter to be A1 2 0(OH) 4 ; (b) for the formation of potash-alum. 6. Why is the tarnish on aluminium the oxide, and not the carbonate (as on Zn and Mg) ? What qualities in a tarnish enable it to protect the metal from further oxidation (p. 423)? 7. Make equations for the action of bicarbonate of soda and aluminium sulphate (alum baking powder) when heated. Explain what raises the bread. 8. Make an equation for the action of: (a) salt on clay at a high temperature; (b) oxygen on indigo white. CHAPTER XXXVII PLANTS, FUELS AND FOOD AGRICULTURE PLANTS and animals are similar in composition. They contain much the same elements, and these are present in the form of similar compounds. They differ sharply, however, in the foods they use in constructing these compounds. Plants use simple, inorganic materials; animals absolutely require complex, organic substances as food. The main chemical processes, therefore, are very different in the two groups. Some of the processes used by plants have been discussed when the substances concerned came up for treatment. We shall here summarize the subject, and complete its outline from the point of view of plant chemistry. This will afford us the basis for understanding the following sec- tions on fertilizers (plant foods), on products of the destructive distillation of vegetable matter, on coal (fossil plants), its uses and products, as well as those on foods, their composition, and diges- tion in the body. PLANTS Plant Composition. We have already learned that plants contain carbohydrates (cellulose, starch, and sugars, Chap. XIX) in large amounts, and esters (vegetable oils, p. 372) in smaller quantities. The proteins are contained very largely in the fruits and also in the protoplasm which is found in every cell. We have already mentioned the gluten (p. 5), a mixture of two proteins, which makes up 10 per cent of wheat flour. In peanuts the percentage of proteins is 26, while in leaves, stems and roots, the moisture is 432 PLANTS 433 very high, and the proteins much lower (lettuce 1.2, potatoes 0.9, turnips 1.3). The proteins are white, amorphous substances which differ from the carbohydrates and esters in that, besides carbon, hydrogen, and oxygen, they contain, on an average, 16 per cent of nitrogen, 1 per cent of sulphur, and in many cases, iron or phosphorus as well. Plants contain, also, compounds of potassium, magnesium, and calcium. Individual plants are noted for containing special compounds. Alkaloids, like quinine from the bark of the cinchona, morphine from the seed capsules of the opium poppy, strychnine from the seeds of nux vomica, are used in medicine. Caffeine gives their stimulating qualities to the extracts from the leaves of tea and the beans of coffee. The laurus camphor a of Formosa is our chief source of camphor. Until recently, twenty-five million dollars worth of indigo was extracted from the Indigofera tinctoria in India alone, although now most of this dye is made from naphthalene (" moth-balls"). There are hundreds of valuable substances each of which is thus found in only one or at most in a very few species of plants. Plant Nutrition. As we watch the growth of a seedling from day to day, we say it "comes up out of the soil." But it does not, entirely the solid part of it comes mostly out of the air. The cellulose, starch, and sugar owe their carbon to carbon dioxide absorbed through the leaves (p. 447). The water (90 to 95 per cent of the whole weight) does indeed come from the soil, as do the nitrogen and sulphur, and the phosphorus and iron required for forming the proteins. The potassium, magnesium, and cal- cium also ascend with the sap. But the carbohydrates are never less in amount than all the other components together (except the water), and are often ten times greater. Fertilizers. The soil is always able to supply the necessary amounts of magnesium, calcium, and iron (as bicarbonates). 434 TEXT BOOK OF CHEMISTRY But it may lack sulphur (as sulphates), nitrogen (absorbed chiefly as nitrates), potassium (as sulphate, chloride or nitrate), and phosphorus (as soluble phosphates). The soil may have been originally deficient in one or more of these necessary plant foods, or the supply may have been exhausted by repeated cropping. Every crop removes permanently a certain part of the supply. Thus, in the case of nitrogen, each crop of maize (corn, 45 bushels) removes 63 pounds per acre, a crop of cabbage (15 tons) removes 100 pounds per acre, clover hay (2 tons) 82 pounds, and wheat (15 bushels) 31 pounds. When the supply becomes reduced, the crops become poor. Moreover, the elements must be in soluble form (p. 433). Felspar and calcium phosphate Ca 3 (PO 4 )2 are not changed by weathering into soluble salts of potash and soluble phosphates with sufficient speed to meet the annual demands (p. 433). It is just as necessary to feed crops as to feed cattle, and equally foolish to starve either of them. Fertilizers are used to make good the original, or acquired deficiences. Potassium salts, chiefly from Stassfurt (p. 358), furnish the potash. Superphos- phate of lime (p. 339) supplies the phosphate. Fertilizers Nitrogen. The nitrogen is applied as sodium nitrate (p. 356), calcium nitrate (p. 320), ammonium sulphate (pp. 260-454), calcium cyanamide (p. 405), manure (pp. 453, 454 or the offal ("tankage") and ground bones from slaughter houses. Most plants probably do not utilize compounds of ammonium, as such. The latter are first oxidized by the help of nitrifying bacteria in the soil (p. 435) to give nitric acid, which, with the lime, gives calcium nitrate. Over every acre of the earth's surface there are 34,000 tons of free nitrogen. Until recently it was supposed that none of this immense store was ever used by plants. It has now been found, however, that peas, beans, clover, alfalfa, and other plants of the order leguminosae, bear round their roots colonies of a kind of FUELS 435 bacteria which has the power to bring free nitrogen into combina- tion. In these root nodules (Fig. 101) the bacteria first produce proteins, which later decompose, and ultimately, by bacterial action, yield nitric acid. In this way a crop of clover will fertilize the soil, not only for itself, but also for the following crop. The advantage of rotation of crops is therefore at length explained. The value of the systematic use of fertilizers is indicated by comparison of the average crop of wheat per acre in different countries. The average of ten successive years is: Denmark 40 bushels, Great Britain 33, Germany 29, United States 14. FIG. 101. FUELS Destructive Distillation of Wood. When dry wood is heated in iron retorts the compounds which it contains (p. 190) are decomposed. Much of the carbon remains in the form of charcoal. The vapors which pass off (through the pipe on the left, Fig. 102) deposit, when cooled, much liquid material. The uncondensed gases are combustible FIG 102. and are used for heating the retorts or other similar purposes. Hard wood furnishes, approximately, 25 per cent of its weight of charcoal, 25 per cent of gases, and 50 per cent of liquids. The liquid contains acetic acid (10 per cent, uses, p. 226), methyl alcohol CH 3 OH or wood spirit (3 per cent, uses, p. 225), a com- plex, tarry mixture, used in road-making (10 per cent), water 436 TEXT BOOK OF CHEMISTRY (77 per cent), and a little acetone (uses, p. 211). The gases con- tain a large part of the nitrogen of the original proteins in the form of ammonia (uses, p. 261), which is dissolved out with water. In the old method, still practised, the wood was stacked, covered with turf (Fig. 103), and set on fire. A part was burned, the rest was turned into charcoal, and all the valuable volatile products were lost. Properties of Wood Charcoal. The charcoal retains the structure a complex network of minute cells of the original wood, and therefore has a surface which is vast in pro- portion to the amount of material it contains. For this reason it will take up many times its own volume of gases, especially of the more condensible ones. Thus, boxwood charcoal takes up am- monia (90 volumes), hydrogen sulphide (55 volumes), and oxygen (9 volumes). For the same reason, when pulverized and shaken with a liquid, it extracts dissolved materials such as dyes (indigo, cochineal, litmus, etc.) and coloringm atters (see sugar refining, p. 222;, especially those of a colloidal nature. Charcoal is used, also, in making gunpowder, in reducing ores, and as a fuel (smokeless). Coal. When wood burns with a plentiful supply of oxygen, it gives nothing but carbon dioxide, water, free nitrogen, and a certain amount of ash (oxides and carbonates of the metals). What happens when it is heated in absence of oxygen, we have just seen. In nature, however, the intermediate case, of slow decomposition of vegetable matter, without much heating and without access of oxygen, takes place on a large scale. Clay and sand, or even simply water, cover the vegetation and exclude the air, and the products are anthracite coal, bituminous coal, or peat. Little is known of the actual compounds contained in coal. We are concerned mainly with the products obtained by heating it in the absence of air, and with its use as a fuel. FUELS 437 Coal Analysis. The moisture is measured by heating 1 g. at 105 for 1 hour. Much water lowers the fuel value, because of the heat wasted in vaporizing it. After reweighing, the sample is heated with a Bunsen flame in a covered crucible to drive off the volatile matter. After weighing again, the air is admitted and strong heat applied to burn up the fixed carbon (coke). The remainder is the ash. For example: Water. Volatile Matter. Fixed C. Ash. Sulphur. Bituminous I . . . . 1.3 36.7 53.5 8.5 1.7 Bituminous II 7 17.9 75.8 5.6 1.2 Anthracite 3.0 5.6 80.5 10.9 0.8 Bituminous coals give much, and widely varying amounts of volatile matter; anthracite coals give very little. The ash is the mineral matter of the original plants, with additional rock materials in some specimens. The coal is selected according to the purpose for which it is to be used. For coal gas, and even for coke, a variety high in volatile matter is chosen. For water gas (p. 203) anthracite or coke itself is employed. FIG. 104. Coal Gas. The gas plant (Fig. 104) includes (1) the fire-brick retorts in which the coal is heated (externally) to 1300, (2) the hydraulic main (a wide iron pipe) immediately above them in which most of the tar collects, (3) the condenser and wash box for cooling, condensing, and removing oils, (4) the scrubbers (vertical and rotary) where the ammonia is taken out by water dripping 438 TEXT BOOK OF CHEMISTRY over strips of wood and by stirring the gas with water, (5) the purifier where hydrogen sulphide is taken up by the hydrated ferric oxide, and (6) the holder in which the gas collects. One short ton (2000 Ibs.) of the bituminous coal I (in above table) gave: Gas 10,500 cubic feet with 13 candle power,* coke 1325 pounds, ammonia 5 pounds (equal to 20 pounds (NH 4 ) 2 S0 4 , worth $60 per ton f), and tar 12 gallons. The percentages of the various components of the gas were: Illuminants (C2H 4 and C 6 H 6 ), 3.8, CO 8.7, H 2 50.5, CH 4 28.2, C 2 H 6 2.8, C0 2 2.1, O 2 0.4, N 2 3.5. Calorific power (see p. 440), 610 B.T.U. per cubic foot. Specific gravity (air = 1) 0.43. By-products. Often the retorts can be opened at both ends, so that the coke may be pushed out by a ram. Frequently they are vertical, so that the charge can be run in, and the coke after- wards removed by gravity. They hold 1000 to 1700 pounds of coal per charge. The ammonia is made into ammonium sulphate. The tar may be used for road-making, as a waterproof material in building, and wherever pitch is applicable. Frequently it is separated by distillation, and other forms of treatment, and yields benzene C 6 H 6 (for aniline and many dyes and drugs), naphthalene Ci H 8 (moth-balls, and starting point for synthetic indigo), anthracene Ci 4 Hio (from which the dye alizarin is made), phenol or carbolic acid C 6 H 5 OH and innumerable other valuable substances. Water gas (p. 203) and carburetted water gas (p. 213) have been discussed already. Coke Ovens. The by-product coke oven is very much like the plant used for making coal gas. The difference is that the * Flame burning 5 cu. ft. per hour. t Garbage (after the extracting of fats, which can be made into soap), when subjected to destructive distillation, yields, amongst other valuable products, no less than 70 to 85 pounds of ammonium sulphate per ton. At present, as a rule, much money is actually spent on throwing these materials away. FUELS 439 heating is arranged so as to decompose the volatile matter and cause it to leave as much as possible of its carbon behind. The resulting gas is consequently poor in illuminants, but excellent as a fuel. The ammonia and tar are also diminished in amount, but are still produced in paying quan- tities. The beehive coke oven (Fig. 105) is a primitive device of fire-brick, shaped like a beehive. It is simply filled with coal, part of which is allowed to burn with a limited supply of air. It yields 66 per cent coke, against 73 per cent from the by-product oven. All the volatile matter, with its gas, ammonia, and tar, escapes through an opening at the top, where it burns in a large flame and is wasted. (For the relative numbers of the two kinds of ovens, see p. 256.) Properties and Uses of Coke. Coke is a grey-black, hard material of spongy texture. It burns without flame, and gives a higher temperature than does coal, because no heat is used in vaporizing moisture and volatile matter. On account of these and other properties, it is used in immense quantities in reducing ores of iron and other metals, and in smaller amounts in electric furnace work and in making electric light carbons. Coal as Fuel. The quality of a fuel coal, and whether it is worth its price, is learned by measuring its calorific (heating) power. A sample (about 1 g.) is burned in a bomb calorimeter. This is a closed, metal vessel, filled with oxygen and submerged in a known weight of water. The coal is set on fire by a wire heated electrically and, after it has burned, the increase in tem- perature of the water is read off. The heat of combustion is the heat in calories (p. 155) evolved by the burning of 1 g. of coal. In engineering practice they use the number of British Thermal 440 Units (1 B.T.U. TEXT BOOK OF CHEMISTRY developed by power. heat required to raise 1 pound of water 1 F.) 1 pound of coal, and call the result the calorific Calories per 1 g. B.T.U. per lib. Calories perlg. B.T.U. per 1 Ib. Hydrogen 28,800 51,840 Coal 7 800 14 040 Charcoal (to CO) 2 .... 8,080 14,544 Anthracite 8,000 14400 Wood (seasoned) 4,750 8,550 Petroleum 11,000 19,800 Knowing that 100 cal. will raise 1 g. of water from C. to 100 C., and 536 cal. more will convert it into steam, it is possible to calculate how much steam should be furnished by 100 kilog. of coal of known heat of combustion. If the quantity falls short, then the furnace, draft, or method of firing may be defective. Too much draft, for example, merely introduces additional, useless material (air) to be heated. Thus, if the flue gas, upon analysis, is found to contain, not 12 per cent carbon dioxide (normal), but only 3 per cent, then for every ton of coal burned, 52 tons of un- necessary air have been raised to the temperature of the furnace. By chemical tests, made in ways like this, the efficiency of every device in the modern factory is (or ought to be) controlled. If the coal is bought without heed to its calorific value, and used without experimental checks, the boiler house alone may easily waste the whole profit earned by the rest of the plant. FOOD Composition of the Human Body. The following girtu* roughly, the percentage of each element in the human body. o 65 N 3 K.. 35 Cl 15 T trace c 18 Ca 2 S .... . 0.25 Mg. 05 F . trace H 10 P .1 Na... . . .0.15 Fe... . .0.004 Si . . . trace We have already learned that the calcium and phosphorus are chiefly in the bones (p. 339). The nitrogen, sulphur, and iron are FOOD 441 in the proteins. The sodium is largely present as salts, in the fluids of the body. The potassium is in the soft tissues and in special secretions like milk. As in the plant, the carbon, hydrogen, and oxygen are in the form of carbohydrates, proteins, and fats, and there is also much water. Elimination. Certain amounts of all these elements leave the system daily. Water evaporates from the lungs and skin. The carbon leaves in large amounts, chiefly from the lungs as carbon dioxide, and also as excreted fats, proteins and carbo- hydrates. Much of the nitrogen is eliminated, chiefly as urea CO(NH 2 )2. The salts are removed in the same way. During the decay of manure, the urea is hydrolyzed, to form carbonic acid and ammonia. The odor of the latter is easily recognizable: CO(NH 2 ) 2 + 2H 2 -> H 2 C0 3 + 2NH 3 . Animal Nutrition. Since there is continual loss, there must be continual replacement. The animal can use, for the most part, only complex compounds (p. 442) . It resembles the plant, in the fact that it can take up into its system only dissolved material. It differs from the plant, however, in the fact that it is provided with a wonderful laboratory in which insoluble substances are changed into soluble ones. This is the digestive tract, consisting of the mouth, stomach, and intestine. The production of soluble sub- stances of suitable composition is called digestion. The processes are too complex for detailed treatment here. Only a few typical illustrations can be given. The principles concerned have all been used and illustrated already, and many of the facts are contained in previous chapters. Foods. First, let us examine the table showing the percentage composition of the edible portion of several articles of food: 442 TEXT BOOK OF CHEMISTRY Food material. Water. Protein. Fat. Carbo- hydrate. Ash. Beef (lean) 73 8 22 1 2 9 1 2 Cod 82 6 15 8 4 1 2 Eggs 73 7 14 8 10 5 1 Milk* 87 3 3 4 5 7 Butter . . ... 11 1 85 3 Cheese (chiddar) . . . 27 4 27 7 36 8 4 1 4 Oatmeal 7 3 16 1 7 2 67 5 1 9 Wheat flour 11.9 13 3 1 5 72.7 0.6 Beans (dried) 12.6 22.5 1.8 59.6 3.5 Almonds 4.8 21.0 54.9 17.3 2.0 Maize (green corn) 75 4 3 1 1 l 19 7 7 Potatoes 78 3 2 2 1 18 4 1 Lettuce . . 94 7 1 2 3 2 9 9 * The emulsified fat separates slowly as the cream; the protein (casein, colloidally suspended in the skim milk) is coagulated by rennet and constitutes cheese; the carbohydrates (lactose, a sugar) is then left in the water, along with inorganic salts. We note, at once, that there is little more water in milk than in cod; that the animal foods, except milk (carrying lactose, p. 442), contain no carbohydrates; that potatoes and corn when dried are nearly all carbohydrate (starch) ; that lean beef when dry is nearly all protein; that some seeds (wheat and beans) contain almost no fat, some (oats) much more, and some (almonds and nuts) a very large amount; and that lettuce and other leaves are mainly water, with dissolved inorganic salts (valuable), contained in a light framework of cellulose (non-digestible) . Digestion of Starch. The carbohydrates, in most foods which contain a large proportion of them, are mainly in the form of starch. The exceptions are milk, sweet fruits, and sugar itself. Starch is insoluble in water, and therefore can not be absorbed. We have seen (p. 220) that, when boiled with a dilute acid, it is hydrolyzed, giving glucose. When bread and potatoes are masti- cated, an enzyme (organic catalyst), named ptyalin, contained in the saliva (alkaline) turns a part of it, by hydrolysis, into a soluble sugar, maltose. Later, in the small intestine, amylopsin completes this process. Here also another enzyme, maltase, splits the mal- FOOD 443 lose into glucose. The glucose then passes through the intestinal wall and so goes into the circulation, where most of it is oxidized. The cooking of starch (baked bread, boiled potatoes, etc.), breaks up the grains and makes the mixing with the enzyme more perfect and the digestion more rapid and complete. Digestion of Fats. Fats, if already emulsified, as in milk, are hydrolyzed by a lipase (enzyme for fat) in the gastric juice of the stomach, and are decomposed into the acid and glycerine (p. 375). Fat in larger masses is hydrolyzed by lipases in the bile and here the acid (insoluble in water) is dissolved. The acid and the gly- cerine then diffuse through the intestinal wall and finally recom- bine to form fat in the blood. Some of this fat is deposited in the tissues and some is oxidized (giving muscular energy and heat). Cooking (application of heat) does not affect the digestibility of fat. However, when fat is heated too strongly, the beginning of destructive distillation produces unsaturated compounds. These are intensely irritating to the digestive organs as the way their vapors bring tears to the eyes would lead us to expect. Digestion of Proteins. The proteins, of which the white of an egg (albumen) is a typical example, are not affected by saliva, but, when mixed with the gastric juice of the stomach, they are changed by the free hydrochloric acid it contains into syntonin. This in turn is hydrolyzed by the pepsin (enzyme), also contained in the gastric juice, into peptones which are soluble in water. These changes, only partly carried out in the stomach, are com- pleted in the small intestine by the trypsin of the pancreatic juice, and the peptones (or amino-acids into which they are split) pass through the intestinal wall into the circulation. The casein of milk, being in colloidal suspension, is completely hydrolyzed to peptones in the stomach. When heated, as in cooking, the proteins do not behave alike. Some, like albumen (white of egg) become coagulated, though 444 TEXT BOOK OF CHEMISTRY probably not less digestible. The same is true of the blood pro- teins (haemoglobin, etc.) of beef. On the other hand, the connec- tive tissue of meat (chiefly collagen) is insoluble in cold water, but in hot water goes into colloidal suspension as gelatine. It is therefore softened by judicious roasting (under-done meat), pro- vided the operation is not carried so far (over-done meat) that the water in the meat is largely evaporated. Fuel Value. While the food is needed to replace the material which is continually eliminated from the system, the organism also requires energy to maintain the routine motions of the heart, intestines, lungs, and other organs, and the normal muscular tension, as well as the movements of the muscles in walking and working. If the heat derived from routine changes is not sufficient to maintain the temperature (37 C.) of the body, then additional material is oxidized by the system for this specific purpose. The values of foods are therefore conveniently estimated in terms of the heat they produce when burned their fuel values. The average fuel values, as measured in the calorimeter, with certain necessary corrections, and expressed, as is usual in this work, in large calories * per gram, are: Carbohydrates 4 Cal., fats 9 Cal., proteins 4 Cal. The fuel values per pound (= 453.6 g.) are 453.6 times greater: Carbohydrates 1800 Cal., fats 4080 Cal., proteins 1800 Cal. Normal Diet. There is much uncertainty, as yet, in regard to the best choice of foods, in respect to the exact distribution in kind and quantity. We know, however, that life cannot be maintained on one kind (say, gelatine) alone. A mixed diet is necessary. In general, it appears that 100 g. of proteins (giving 4 X 100 Cal.) per day, and a sufficient amount of other foods to bring the total fuel value up to 2200 Cal. per day, is sufficient for a * One large calorie (1 Cal.) is equal to one thousand small calories (1000 cal.), as hitherto denned (p. 155) and used. FOOD 445 person leading a strictly sedentary life. For work involving physi- cal exercise, larger values, up to about 3800 Cal., are required. From the data given in the table (p. 442) the fuel value of 100 g. of each kind of food can easily be calculated. Fuel Values and Prices of Foods. If the current prices are considered, one can also readily calculate the fuel value obtainable for a given sum of money invested in each kind of food. Thus: lean beefsteak contains 22.1 per cent of protein, or 0.221 pounds per pound of meat. The fuel value of this protein is 0.221 X 1800, or 398 Cal. per pound. Price* per 1 Pound, Cents. Fuel Values per 1 Pound Cal. per 10 Cents. Protein. Fat. Carbohyd. Total. Steak 25 24 5 5 40 2.5 25 398 266 290 249 378 40 500 118 428 294 61 2240 4 1500 516 694 1799 1619 2929 375 2074 206 290 3600 3240 732 1500 830 Eggs. . Oatmeal 1215 1309 311 331 74 Flour Almonds Potatoes Cheese * The prices vary greatly with the quality, the season of the year, the demand, the supply, etc. Animal Charcoal. When bones or dried blood are subjected to destructive distillation, gases and vapors are given off and charcoal remains. The charcoal from bones (bone black) contains 90 per cent of mineral matter (largely Ca 3 (PO4)2) and only 10 per cent of carbon. Both kinds of animal charcoal, being very active as decolorizers, are used in sugar refining. Exercises. 1. Give the uses of: (a) acetic acid; (b) methyl alcohol; (c) acetone; (d) each form of carbon mentioned in this chapter or elsewhere in the book. 2. (a) Why are charcoal and coke smokeless fuels? (b) Ex- plain why bituminous coal burns with flames while anthracite does not. 446 TEXT BOOK OF CHEMISTRY 3. In making charcoal, why is the turf used? 4. Point out the analogies between the processes used in making coke and charcoal, and between their properties and uses. 5. What gas is given off when vegetable matter decomposes under water? 6. What compounds contained in the original plants, furnish: (a) the hydrogen sulphide; (b) the ammonia obtained on distilling coal? 7. A gas of sp. gr. 0.43 (air = 1) gives, on burning, 610 B.T.U. per cu. ft. How many B.T.U. is this per pound? 8. How many kilog. of steam can be made by burning 100 kilog. of coal, the heat of combustion of which is 8500 cal. ? 9. How many pounds weight of steam can be made by burning 50 Ibs. of coal of calorific power 14,500 B.T.U. 10. At 5 atmos. pressure and 152 C., how many cubic meters will 190 kilog. of steam occupy? 11. Find the shortest method to prove that the calorific power (p. 440) is always numerically 1.8 times the heat of com- bustion of one gram. 12. What is " conservation "? What four industries or opera- tions (or ways of performing operations) that are wasteful have been mentioned in this chapter (compare p. 440) ? 13. Rewrite the statement on p. 445, lines 11-19, in terms of B.T.U., degrees F. and pounds. 14. By what changes does stable manure furnish food for plants? 15. Make equations for the formation: (a) of maltose from starch; (b) of glucose from maltose. 16. Make a connected statement showing the stages in the digestion of milk. 17. If we regard the haemoglobin in the red blood-corpuscles as a catalyst, what actions does it catalyze? 18. Why does fat appropriately form a larger proportion of the diet in the Arctic regions than elsewhere ? AGRICULTURE 447 19. Give the weights of carbohydrate, protein, and fat which would supply a menu, such that the total food value was 3000 Cal., and that 100 g. of protein was included, and the remaining fuel value was. divided equally between carbohydrates and fats. 20. Calculate the calorific value of 1 kilog. of: (a) wheat flour, (b) oatmeal. AGRICULTURE * Plant Food. For successful growth, plants require carbon, hydrogen, oxygen, phosphorus, potassium, nitrogen, sulphur, calcium, iron, and manganese.t The carbon is chiefly supplied by the carbon dioxide in the air, which enters the leaves through the stomates. Water supplies the hydrogen. Oxygen enters through the leaves, and also through the roots and stems. The oxygen produces oxidation of substances in the plant, and gives heat. Phosphorus, nitrogen, and sulphur are absorbed by the roots as soluble phosphates, nitrates, and sulphates, and nitrogen sometimes as ammonia. Potassium commonly enters as carbon- ate or bicarbonate. Calcium and magnesium are absorbed as phosphate, nitrate, sulphate, or bicarbonate; and iron as ferric hydroxide (Fe(OH) 3 ). All plant food must be soluble. A weak solution is even better for a plant than is a strong solution. Even ferric hydroxide dissolves sufficiently in soil moisture to furnish the necessary supply of this substance. The solutions in the soil are absorbed by the minute root-hairs, which are near the tips of the roots. The protoplasm which lines the root-hairs exercises a selective action, permitting the entrance of certain substances, and the exclusion of others. The protoplasm layer also prevents * This chapter contains only a brief outline of substances used by plants as food, and of the conditions which favor the growth of plants. As soils in different regions vary considerably, the instructor will have to supplement this chapter by giving a more thorough discussion of the soil hi the neighbor- hood where he is teaching. t Manganese and chlorine are also used by many plants. In some cases, sodium salts can take the place of potassium salts. 448 TEXT BOOK OF CHEMISTRY the loss of sugars and of soluble proteins. The roots also keep the plant upright. Formation of Sugar and Starch. Only about one to two per cent of sunlight is used the rays which are effective being in the orange-red part of the spectrum. It is the chlorophyl which absorbs these rays and enables them to be used. Sugar and starch are formed from carbon dioxide and water. The former enters the leaves through the stomata, the water enters through the roots and meets the carbon dioxide in the leaves. The carbonic acid changes into formaldehyde and oxygen: H 2 CO 3 -+CH 2 + 2 . The oxygen escapes through the stomata. The formaldehyde, dissolved in water, quickly changes to dextrose: 6CH 2 0-C 6 H 12 6 . The dextrose is affected by an enzyme, which removes water, and starch is formed : n C 6 H 12 6 - n H 2 -> (C 6 Hi 5 )n. These changes take place only in daylight. One acre of corn, when the ears are developing, will produce about 170 pounds of carbohydrates in one day. Dextrose also changes to cellulose, which forms the hard casing of the stem of a plant. The starch first accumulates in the leaves, but is later turned back into dextrose (soluble) by amylases, and then passes from cell to cell until it enters the sieve tubes. It then moves to parts which require new material, such as growing leaves, roots, flowers, or growing tips. Starch is insoluble. Some oxidation takes place in a plant, and some carbon dioxide is liberated, by oxidation of carbohydrates. Oils and Proteins. Particles or films of oil are found in plant cells. They seem to be made by reduction of dextrose. AGRICULTURE 449 Proteins are the most complex substances a plant contains. The proteins contain carbon, hydrogen, oxygen, and also nitrogen, sulphur, and frequently phosphorus. The proteins are formed during daylight. Protoplasm. Protoplasm is the real seat of life in the plant. It is a soft, slimy material suspended in water. The approximate composition of dried protoplasm is: 50-66 per cent proteins, 25-17 per cent oils and carbohydrates, and 25-17 per cent of organic acids, organic bases, and mineral salts of potassium, magnesium, calcium, and iron. Seeds. Seeds contain carbohydrates, oils, and proteins, and contain little water. Some plants produce large fleshy roots, like the beet, or underground tubers like the potato, or enlarged bulbs like the onion. The potato is composed mainly of starch. Phosphorus. Phosphorus is a necessary constituent of some proteins. Cell nuclei contain phosphorus, which stimulates cell division, and promotes growth. Soluble phosphates also stimu- late the growth of seeds, and advances the ripening of seeds. Phosphates are also useful in the ripening of grain, and likewise stimulate the development of root growth, which is of great importance in dry weather, Potassium. It seems likely that organic compounds, con- taining potassium, assist in the change of formaldehyde to dex- trose. Potassium salts also assist in the decomposition of starch into soluble sugars. Potassium also plays a part in the formation of cellulose. Potassium salts also assist the plant to resist the growth of fungi, and assist materially in giving a good potato crop. Potassium also increases the yield of sugar in beets. Nitrogen and Sulphur. Nitrogen is contained in all pro- teins. It assists in the growth of new stems and leaves, it is valuable in the growth of lettuce, cabbage, and celery. 450 TEXT BOOK OF CHEMISTRY Sulphur is contained in all proteins, and produces flavoring oils in mustard, onion, cabbage, and horseradish. Calcium, Iron, Magnesium. Calcium assists in the solu- tion and transportation of starch, it assists in the production of strong cell walls, and numerous root-hairs, it precipitates oxalic acid, as calcium oxalate, for oxalic acid is a poison to plants. Iron is required for the formation of chlorophyl, although not contained in chlorophyl. Magnesium salts assist in the formation of chlorophyl, they take part in producing proteins containing phosphorus, and they assist in the formation of seeds. Alkaloids in Plants. Atropin is a poisonous substance, and is used to dilate the pupil of the eye. Caffeine and theine are found in coffee and tea, and make coffee and tea stimulants. Cocaine is another poisonous substance. Morphine is a valuable narcotic, but it is apt to lead to continuous use, and produces harmful effects both on mind and body, when used frequently. Nicotine is found in tobacco, from one to eight per cent. When tobacco is burned, the nicotine is decomposed, but when tobacco is chewed, this poison enters the system. Bacteria. One gram of soil contains 100,000 bacteria, and sometimes 100,000,000 bacteria. The bacteria assist in the de- composition of vegetable or animal matter in the soil. Bacteria assist in the decomposition of mineral particles, in the formation of nitrates and other soluble inorganic salts, which serve as plant foods, and, by producing humus, increase the moisture-holding capacity of the soil. Nitrification of the Soil. Bacteria generally liberate the nitrogen from organic matter as ammonia (NH 3 ). Other organ- isms assist in the oxidation of ammonia to nitrous acid: 2NH 3 + 30 2 -> 2HNO 2 + 2H 2 0. AGRICULTURE 451 Later the nitrous acid is oxidized to nitric acid with the assistance of nitrate bacteria : 2 -+2HN0 3 . Another method by which nitrogen is obtained from the air is by the action of certain bacteria surrounding the roots of clover, alfalfa, peas, beans and other legumes. Proteins are produced in root-nodules, nitrogen from the air being used, and then after- wards decompose giving soluble compounds containing nitrogen. Substances (bacteroids) produced by bacteria enter the roots, and stimulate growth.* Source of Soluble Phosphates. Apatite Ca 5 (P04) 3 Cl (or F) furnishes soluble phosphates. Under the action of water and carbon dioxide, it changes as follows: 2Ca 5 (P0 4 ) 3 Cl + 12C0 2 + 12H 2 O ** SCaH^PO^ +- 6CaH 2 (CO 3 ) 2 + CaCl 2 . Calcium phosphate reacts in the same way: Ca 3 (PO 4 ) 2 + 4CO 2 + 4H 2 - CaH 4 (PO 4 ) 2 + 2CaH 2 (C0 3 ) 2 . Monocalcium phosphate CaH 4 (PO 4 ) 2 is soluble to the extent of 1 part in 100 of water. Potassium Salts 9 Sulphur and Calcium Minerals. Solu- ble salts of potassium are usually obtained from feldspar K 2 Al 2 (Si 3 Os) 2 . Water containing carbon dioxide decomposes this mineral: K 2 Al 2 (Si 3 8 ) 2 + C0 2 + 2H 2 -> Al 2 (OH) 4 Si 2 5 + 4Si0 2 + K 2 C0 3 . The potassium carbonate supplies the potash needed by plants. Calcium sulphate CaS0 4 ,2H 2 is sufficiently soluble to supply sulphur to plants. * Azobacteria have the power to use atmospheric nitrogen, and produce nitrogen compounds, although the amount produced is not very great. 452 TEXT BOOK OF CHEMISTRY Calcium carbonate is a common mineral. When it comes in contact with water containing carbon dioxide, it becomes more soluble : CaC0 3 + H 2 + C0 2 - CaH 2 (C0 3 ) 2 . The dissolved salts in the water contained in the soil adhere to the minute particles of the soil, and are not readily washed away by rain. This adhesion is called adsorption. Fertilizers Nitrogen Compounds. Planting clover or alfalfa will increase greatly the amount of available nitrogen in the soil. Sodium nitrate can also be spread on the soil. Also calcium nitrate Ca(N0 3 ) 2 , made in Norway, can be employed. The latter salt is better than sodium nitrate. It assists in making potassium salts available, and also improves the physical condition of the soil. Ammonium sulphate (NH 4 ) 2 SO 4 is also employed in con- siderable amounts. This salt, if used continuously, produces an acid condition in the soil, which interferes with the growth of some crops. Cyanamide is made from calcium carbide: 3C-CaC 2 + CO. The calcium carbide is then treated with nitrogen at a high tem- perature, giving cyanamide: The cyanamide is hydrolyzed by water in the soil: CaCN 2 + 2H 2 - H 2 CN 2 -f Ca(OH) 2 , H 2 CN 2 + H 2 O '-+ CO(NH 2 ) 2 (urea), CO(NH 2 ) 2 + 2H 2 - (NH 4 ) 2 C0 3 (Am. carbonate). The ammonium carbonate is first converted into nitrous acid and nitrites, and then into nitric acid, and nitrates. The calcium hydroxide neutralizes the nitric acid, giving a basic calcium nitrate. Refuse from packing houses, called tankage, is often used as a fertilizer. After the oil has been pressed out of cotton seed, the residue may be used as a fertilizer. AGRICULTURE 453 Sources of Phosphates. When bones are heated strongly in a closed vessel, the organic matter is decomposed, and Ca 3 (PO 4 ) 2 with ten per cent of carbon remains. When this material is treated with sulphuric acid, we get a soluble phosphate: Ca 3 (PO 4 ) 2 + 2H 2 S0 4 -> CaH 4 (P0 4 ) 2 + 2CaS0 4 . Large supplies of calcium phosphate are now available in Florida, South Carolina, Tennessee, Idaho, Wyoming and Montana. When treated with sulphuric acid, as above, and mixed with manure, it forms an excellent fertilizer. Basic slag or Thomas slag, a phosphate of calcium, obtained from British blast furnaces, is also used as a f ertilizer. Phosphates, containing much iron, should be avoided. Potassium Fertilizers. Much potassium chloride is ob- tained from Strassfurt in Germany. Potassium sulphate is also used. It has no bad effect on tobacco or potatoes, but does make the soil acid. Wood ashes contain much potassium carbonate, and are used as fertilizers. Lime. Alfalfa uses 212 pounds of CaO per acre (600 pounds per acre are needed, as 460 pounds are washed way), cabbage uses 95 pounds per acre, clover uses 90 pounds per acre, tobacco uses 83 pounds per acre. So that a good supply of lime is needed for these and some other crops. Sandy soil also requires CaO also. It can be applied as calcium carbonate (limestone) or as calcium hydroxide. Ground limestone is also used, immediately after plowing, to neutralize a soil which contains much clay, as clay is an acid salt, H 2 Al 2 (SiO 4 ) 2 . Manure. The solid excrement from animals contains undi- gested food. The proteins suffer some decomposition in the stomach, and further decomposition due to bacteria in the in- testines. The bacteria are valuable in promoting decomposition in the soil.- Further decomposition is needed, before the material becomes soluble and available for plants. Urine of animals also 454 TEXT BOOK OF CHEMISTRY furnishes soluble food for plants. The excrements contain, on an average, about eighty per cent of the food eaten. One ton of farm manure contains about 10 pounds of nitrogen, 5 pounds of phosphoric acid, and 10 pounds of potash. The manure is mixed with the soil just before planting seed, or is used as top dressing. The air assists materially in the changes in the manure. Thus urea is hydrolyzed: CO(NH 2 ) 2 + 2H 2 O -> (NH 4 ) 2 C0 3 -> 2NH 3 + C0 2 + H 2 In presence of sufficient water, the ammonium carbonate does not decompose. The proteins are changed by the air into ammonia. The formation of nitrates seldom happens in manure piles, but scattered manure develops nitrates. The potassium forms potassium hydrogen carbonate KHC0 3 . The phosphorus and sulphur become soluble phosphates and sul- phates. Molds assist in the chemical changes. Often half of the nitrogen, phosphorus, and potassium are washed away. To prevent loss of nitrogen, gypsum CaSO 4 ,2H 2 O is added to manure at the rate of 100 pounds per ton. This gives ammonium sulphate : (NH 4 ) 2 C1 + CaS0 4 -* (NH 4 ) 2 S0 4 + CaCO 3 . The calcium carbonate assists in neutralizing acid soils. Lime should never be used, as it liberates the ammonia as a gas. Wood ashes (potassium carbonate) will drive off ammonia more readily than lime. Also, acid phosphate (40-50 pounds to each ton of manure), or rock phosphate (80-100 pounds to each ton of ma- nure). Commercial fertilizers kill seeds, so they have to be mixed with the soil just before the seeds are planted. Works of Reference. Edward J. Russell, Soil Conditions and Plant Growth (Longmans, Green & Co.). Charles W. Stod- dart, The Chemistry of Agriculture for Students and Farmers (Lea and Febiger). Lucius L. van Slyke, Fertilizers and Crops, or the Science and Practice of Plant-feeding (N. Y. Orange Judd Co.; London, Kegan Paul, Trench Trubner & Co.). CHAPTER XXXVIII IRON, NICKEL, COBALT HAVING disposed of most of those metallic elements, the chemis- try of which, because of the fact that each has only one valence, is comparatively simple, we turn now to the metal which is by far the most important of all in the industrial standpoint, namely iron. IRON FE Occurrence. Particles of free iron are found in igneous rocks, and larger pieces in meteorites. Compounds of iron are very widely distributed, in the soil and rocks (giving yellow and red tints), in the chlorophyl and proteins of plants and in the haemo- globin and other proteins of animals. Pyrite FeS2 (fools' gold) is used mainly as a source of sulphur for sulphuric acid. The ores which yield iron itself are: (ferric oxide), red haematite. Red when pulverized. 2Fe203,3H 2 (hydrated ferric oxide), brown haematite. FesOi (magnetic oxide of iron), magnetite. Black when pul- verized. FeCOs (ferrous carbonate), spathic iron ore. The carbonate, mixed with clay (clay iron-stone), furnishes most of the iron in Great Britain, but less than one per cent of it in the United States. The ore is first calcined to produce the oxide. The Blast Furnace. Coke is used to reduce the oxides and as fuel. The carbon monoxide CO, produced by the burning of the coke and air, is the actual reducing agent : Fe 2 3 + 3CO^2Fe(melted) + 3C0 2 T 455 456 TEXT BOOK OF CHEMISTRY Since, however, the ores contain rocky material (gangue), such as silica SiC>2 and silicates of aluminium, limestone is added in the proportion required to give a fusible slag. Thus the whole material is converted into gases which ascend, and liquids (melted iron and slag) which flow down. The blast furnace (Fig. 106) is an iron structure 40 to 100 feet high, lined with fire-brick. A circular pipe delivers compressed air to several nozzles (tuyeres). The ore, coke and limestone are admitted from the hopper at the top by lowering two bells, only one of which is opened at a time. The gases escape at openings below the bells into the down- comers. The highest temperatures are reached, and the most chemical action occurs, at the widest part. The melted iron and slag (immis- cible) collect in two layers in the hearth or crucible at the bottom. From time to time the slag is allowed to flow from an opening near the top of the crucible, and the iron from a similar opening at the bottom and 90 from the former. Plugs of wet clay close the openings and are in- stantly baked hard. The iron is taken in ladles to other parts of the plant, or is cast into "pigs" in steel moulds and chilled in water. When the proportions of the material are properly controlled by analyses, the furnace runs continuously until the lining is worn out. Reactions in the Blast Furnace. The action (see equation, above) is a reversible one and the carbon dioxide formed reacts with the iron to reproduce the original materials. The system reaches equilibrium when about 1 volume of carbon dioxide is present with 1 J volumes of carbon monoxide. To keep the action going forward there must, therefore, be constantly present an excess of carbon monoxide beyond this amount, namely about 2 FIG. 106. IRON Fe 457 vols. CO : 1 vol. CO2. This means that the gases issuing at the top contain the same proportions, or about 20 per cent CO, 10 per cent C0 2 , 70 per cent N 2 . The downcomer gas is therefore combustible, and is used in the blowing engines, and for heating the air-blast to 400 to 550. From 3 to 4 tons of air are required for every ton of iron produced. The moisture in this air acts upon the coke, giving water gas (p. 203). This uses up fuel, and also lowers the temperature just where it ought to be highest. In" the most modern furnaces, therefore, the air blast is regulated in accordance with the amount of moisture the air contains. In 1911, 23,250,000 tons of iron were produced in the United States. Cast Iron. Pig iron contains 4 to 5 per cent of carbon and varying amounts of silicon (as silicide of iron), phosphorus (as phosphide) and sulphur (as sulphide). These impurities lower the melting-point from 1510 to about 1100. The material is hard and brittle. Most of it is made into wrought iron or steel, but some is used in making objects of cast iron, such as ranges, stoves, pipes, and radiators, which are not to be subjected to shocks or strains. Cast iron expands on solidifying and forces itself into the details of the mould. By adding pyrolusite Mn0 2 in the blast furnace, cast iron con- taining from 20 per cent of manganese (spiegeleisen) up to 80 per cent (ferromanganese), and carbon up to 6 per cent, is made for use in steel manufacture. Wrought Iron. Wrought iron is commercially pure iron. The broken pigs are placed in a reverberatory furnace (Fig. 107), the hearth of which is covered with a bed of haematite ore Fe 2 O 3 and silicates. The flames and heated gases, deflected by the low roof, play upon the iron and melt it. The oxygen in the haematite combines with the carbon, phosphorus, sulphur, and silicon, giving the oxides. The mass is worked vigorously with iron rods upon 458 TEXT BOOK OF CHEMISTRY the bed of haematite (puddled), carbon monoxide escapes, and the iron becomes more viscous as its melting-point rises on account of the removal of the impurities. Finally, it is collected in balls (blooms) on the iron rods. The treatment occupies an hour and a half. To press out the slag, the blooms are first passed through a squeezer and then put through the rolls. The result- FIG. 107. ing bars are repeatedly cut, " piled" in a bundle, reheated, and rolled. This treatment, and the presence of a little slag dis- tributed through the mass, give wrought iron the valuable prop- erties which distinguish it from all other iron products, namely its fibrous fracture and its extreme toughness. On account of these properties it is used for anchors, chains, and bolts. It is drawn into wire, and, when heated, can be cut, shaped, and welded under the hammer. The impurities having been greatly reduced (to 0.1 or 0.2 per cent), this iron is much less fusible than cast iron, and is used for fire bars. Crucible Steel. Steel contains 0.75 to 1.5 per cent of carbon, and is freed as far as possible from other impurities. Small lots, for special purposes, are made in clay (or graphite and clay) crucibles in melts of 60 to 100 pounds. The charge in Sheffield consists of blister steel, i.e., carburized Swedish wrought iron of varying carbon content. The modern method is to melt " melt- ing bar/' a very pure open hearth steel with charcoal, or even IRON Fe 459 pure pig iron. Crucible steel is used in making razors (1.5 per cent C), tools (1 per cent C), dies (0.75 per cent C), pens, needles, and cutlery. Electric heating (e.g. in the Heroult furnace), recently intro- duced, permits the steel maker, first to wash the molten iron with basic slags of high oxidizing power until perfectly pure, and then by suitable additions to give it any required final composition. Properties of Steel. Cast iron can be melted and cast, is hard when chilled, but can not be forged or rolled. Wrought iron is slag-bearing and malkable, and is not hardened by quenching from a high temperature. It is never cast. Steel is free from slag, being cast from an originally liquid condition. If its carbon content is high enough, it can be hardened by quenching. Steel has also greater tensile strength * than wrought iron, which in- creases with the carbon content, and it can be permanently mag- netized. In addition, high carbon steel can also be tempered (p. 459) to the required degree of hardness. tempering. To understand the last fact, it must be noted that the carbon in steel is in the form of carbon or of a carbide of iron FesC (6.6 per cent C) homogeneously dissolved in the iron. When white hot steel (up to 2 per cent C) is suddenly chilled, there is no time for any changes to occur during the cooling, and a solid solution is obtained which is very hard and brittle. When, however, the cooling is slow, some of the carbon separates in minute crystals of cementite, carbide of iron Fe 3 C, until at about 700 there remains only 0.9 per cent of carbon in solution. At this temperature the solid solution breaks down into a mechanical mixture of pure iron which is soft, and carbide of iron which is hard. Steel is thus a mixture, and not homogeneous, when slowly cooled. When therefore hard, chilled steel is heated once * Tensile strength or tenacity is measured by the weight (in kilograms) required to break a wire of the metal 1 sq. mm. in section. Lead 2.6, copper 51, iron 71, steel 91. 460 TEXT BOOK OF CHEMISTRY more for the purpose of tempering, the extent to which the softer material is formed depends on the temperature reached, and on the rate and duration of the cooling permitted. By varying these the degree of hardness allowed to remain can be adjusted. Phosphide of iron makes steel brittle when cold ("cold short"). Sulphide of iron makes it brittle when hot ("red short"), and unsuitable for forging. Hence phosphorus and sulphur are re- duced to the lowest possible amounts. Bessemer Process. Pig iron is melted and run into the converter (Fig. 108), which is lined with firebrick, measures about 15 by 8 feet, and holds 15 tons. An air-blast, entering through one axle, blows through the metal from tuyeres at the bottom. The oxidation of the carbon and silicon, which raises the temperature above the melting-point of pure iron, is over in 20 minutes. Spiegeleisen IG ' ' is then added to give the desired percentage of carbon and manganese, and the liquids, first the slag and then the iron, are poured into ladles, and the metal is cast. Sulphur and phosphorus are not removed by the air. If present in too great amounts, they are removed by lining the converter with basic material such as magnesium and calcium carbonates (Thomas-Gilchrist process). The slag then contains phosphates, and is valuable as a fertilizer. Bessemer steel is used for heavy and light machinery castings, and is rolled into bridge and structural iron. It contains from 0.1 per cent (soft) to 1 per cent (hard) of carbon. Open-hearth (Siemens- Mar tin) Process. In this process pig iron and scrap iron (up to 75 tons) are melted on a hearth lined with firebrick and sand (Fig. 109), At a later stage haematite is IRON F6 461 added to furnish oxygen (as in puddling) . To secure economically the temperature necessary to keep the pure product (iron) fused, Siemens contributed the idea of preheating the fuel gas and air by a regenerative device. The spent air and gas pass out through a checkerwork of brick. When this has become hot, the valves are reversed, the gas and air now enter through the hot brickwork FIG. 109. and pass out through the checkerwork on the opposite side, raising its temperature. The direction of the gases is changed every twenty minutes or so, and the whole operation lasts 8 to 12 hours. Towards the end some aluminium is added to combine with oxygen (present as CO), and give sounder ingots. Recently, iron containing 10 to 15 per cent of titanium has been added instead. The titanium combines with both nitrogen and oxygen, and passes into the slag. Rails made of steel purified with this element are less liable to breakage and are 40 per cent more durable than ordinary open-hearth rails. There are nearly a million tons of these rails now (1912) in use in the United States. 462 TEXT BOOK OF CHEMISTRY The advantage of this process over that of Bessemer is that it is not hurried and is therefore under better control. The material can be tested by sample at intervals, and defects cor- rected. The product is of better and more uniform quality. As in the Bessemer process, phosphorus and sulphur are re- moved by using a basic lining. Open-hearth steel is used for the better class of rails, for railway bridges, for shafts, armor-plate, and heavy guns, and wherever the steel is subject to much vibration. Steel Alloys. We must distinguish between manganese, aluminium, silicon, or titanium added in small amounts (''med- icine") to purify the iron, and passing (in combination) into the slag, as described in preceding sections, and the present sub- ject, which concerns metals added so as to produce regular alloys. Manganese steel (7 to 20 per cent Mn) is exceedingly hard even when cooled slowly. It therefore does not lose its temper readily when heated by friction. It is used for the jaws of rock-crushing machinery and for burglar-proof safes. Chromium- vanadium steel (1 per cent Cr, 0.15 per cent Va) has great tensile strength, can be bent double while cold, and offers great resistance to changes of stress, and to torsion. It is used for frames and axles of automobiles, and for connecting rods. Tungsten steel (tungsten 8 to 20 per cent, and chromium 3 to 5 per cent) is used for tools in high-speed metal turning, as it can become red hot (from friction) without loss of temper. Its use has greatly increased the ratio of output to wages. Nickel steel (containing 2 to 4 per cent of nickel) resists cor- rosion, and has a very high limit of elasticity and great hardness. It is used for armor-plate, wire cables, and propeller shafts. The alloy with 36 per cent nickel, called invar, is practically non- expansive when heated and is valuable for meter-scales and pendulum rods. IRON re 463 Properties of Pure Iron. Pure iron may be made by electrolysis, or by reduction of a pure salt by hydrogen. It has a white luster, is very tough, and does not melt below 1510. The purest iron does not rust in cold water. Ordinary iron rusts in moist air or under water, forming a hy- drated ferric oxide 3Fe 2 O3,H 2 0. The impurities act as contact agents. The rust is a brittle, porous, non-adherent coating, which does not protect the metal below. Oil protects iron from rusting because, although oxygen penetrates the oil, being soluble in it, moisture does not. Iron displaces hydrogen from hydrochloric and sulphuric acids, giving ferrous salts : Fe + 2HC1 -> FeCl 2 + H 2 t . The impurities carbide, sulphide, and phosphide produce hydrocarbons, hydrogen sulphide, and phosphine PH 3 , and the last two confer an odor on the escaping gas. Iron burns in oxygen, and acts when heated upon steam, in both cases producing magnetic oxide of iron Fe 3 4 (p. 51). A thin film of this oxide is adherent, and protects the iron ("Russia" iron). Iron has Two Valences. One atomic weight of iron (55.84) may hold two or three atomic weights of a univalent element in combination. Thus, we have ferrous chloride FeCl 2 and ferric chloride FeCl 3 , the bromides FeBr 2 and FeBr 3 , the oxides FeO and Fe 2 3 , ferrous sulphate FeS0 4 and ferric sulphate Fe 2 (SO 4 ) 3 , and so forth. It may therefore be bivalent or trivalent, according to the chemical conditions. For example, in the action of iron on a dilute acid (see equation, above), no ferric salt is formed. Indeed, if ferric chloride is added, the hydrogen and iron reduce it to ferrous chloride (see p. 464). The ferrous salts are pale green and give colorless solutions, containing ferrous-ion Fe" 1 ^. The ferric salts, containing the ion Fe 4 " 1 ^, are usually yellow in solution, on account of ferric hydrox- ide produced by hydrolysis. The differing properties of the two 464 TEXT BOOK OF CHEMISTRY ions, and the change from ferrous to ferric, and vice versa, are dis- cussed below. Ferrous Sulphate FeSO 4 . When the bath of dilute sul- phuric acid, used in cleaning iron for making tin-plate (p. 478). and galvanized iron (p. 419), is becoming exhausted, scrap iron is thrown in to use up the rest of the acid. The solution gives, on evaporation, pale green crystals of ferrous sulphate, FeSO4,7H 2 (copperas or green vitriol). The salt is used in making ink (see p. 467) and rouge (see p. 465), and in purifying water (p. 427). Chlorides of Iron. Ferrous chloride FeCl 2 is obtained in solution when iron displaces hydrogen from hydrochloric acid (p. 52), and is isolated by evaporation. The hydrate FeCl 2 ,4H 2 is pale green, the anhydrous salt colorless. When chlorine is dis- solved in the solution, or when the latter, acidified with hydro- chloric acid, is exposed to the air, ferric chloride FeCl 3 is produced : 2FeCl 2 + Cl 2 -*2FeCl 3 or 2Fe++ + C1 2 -> 2Fe+++ + 2C1~ 4FeCl 2 + 2 + 4HC1 - 4FeCl 3 + 2H 2 O and is familiar in the form of a yellow hydrate FeCl 3 ,6H 2 ob- tained by evaporation. Other oxidizing agents, such as nitric acid, produce the same change. Ferric chloride, in solution, has an acid reaction, due to hydroly- sis. It is reduced to ferrous chloride by shaking the solution, or more quickly by boiling it, with iron filings. Fe-+3FeCl 2 or 2 Fe+++ + Fe -> 3Fe++. Hydroxides of Iron. Ferric hydroxide Fe(OH) 3 appears as a brown precipitate when an equivalent amount of sodium hy- droxide is added to a solution of a ferric salt: FeCl 3 + 3NaOH - Fe(OH) 3 J, + 3NaCl. When only a little sodium hydroxide is added, the brown pre- cipitate, formed locally, redissolves to give a deep reddish solution. IRON Fe 465 This contains ferric hydroxide in colloidal suspension. The sodium chloride and unused ferric chloride can be separated by dialysis (p. 378), and a pure colloidal suspension of the hydroxide obtained. Ferrous hydroxide Fe(OH) 2 produced by precipitation, is white when pure, but becomes quickly green and then brown by oxida- tion. Oxides of Iron. When ferric hydroxide is heated, ferric oxide Fe2O 3 remains as a red mass: 2Fe(OH) 3 ->Fe 2 O 3 + 3H 2 O f . It is made by calcining (roasting) ferrous carbonate, ferrous sul- phide, or ferrous sulphate: 4FeSO 4 + O 2 -> 2Fe 2 3 + 4SO 3 . 4FeS + 70 2 -* 2Fe 2 O 3 + 4SO 2 . The oxide, when pulverized in a ball mill, gives a powder of more or less brilliant red color commonly used in paints (Venetian red and Indian red). That from ferrous sulphate is rouge, used in polishing plate glass and lenses, and as a pigment. Yellow ochre is a natural hydrated ferric oxide 2Fe 2 O 3 ,3H 2 O, which acquires various depths of color during calcination, and constitutes the sienna and umber used for paints. Magnetic Oxide of Iron Fe 3 O 4 is regarded as a compound of ferrous and ferric oxides FeO,Fe 2 3 . It is formed by strongly heating ferric oxide: 6Fe 2 O 3 ->4Fe 3 4 + 2 and is formed when iron is oxidized at a high temperature by oxygen, air (blacksmith's scale), or steam. It can be magnetized, and natural specimens are often strongly magnetic (lodestone). Ferrous oxide FeO is a black substance made by cautious re- duction of ferric oxide by a stream of hydrogen. 466 TEXT BOOK OF CHEMISTRY Ferrous Carbonate FeCO 3 . The carbonate occurs in nature as an impurity in clay, in clay iron stone, and pure as siderite. Water containing carbonic acid dissolves it, giving the bicarbonate : FeC0 3 + H 2 C0 3 -> Fe(HCO 3 ) 2 . Thus well and river waters all contain at least traces of ferrous bicarbonate as a part of their hardness. Exposure to the air causes oxidation and, as ferric carbonate is not stable, rust (ferric hydroxide or a hydrated ferric oxide) is deposited: 4Fe(HCO 3 ) 2 + 2 + 10H 2 0-+4Fe(OH) 3 + 8H 2 C0 3 . This red deposit is seen in white vessels in which such water drips or stands. It also " yellows" goods washed in such water, if the carbonate of iron is not previously precipitated (p. 402) by soda or some other softening agent (see p. 403). Ferro- and Ferri- cyanides. Potassium ferrocyanide K 4 Fe- (CN) 6 , or yellow prussiate of potash, is a pale yellow, soluble salt. The iron is contained in the negative radical and ion Fe(CN) 6 == , and the solution therefore gives the reactions of this ion, and not of ferrous- or ferric-ion. One of the double decompositions of this salt namely, that with ferric salts is important because it gives a gelatinous precipitate of Prussian blue (ferric ferrocyanide) : 3K 4 Fe(CN) 6 -Fe 4 m (Fe(CN) 6 ) 3 IV J, + 12KC1. Prussian blue is employed in making paints, and is the usual pigment in laundry blueing. Although insoluble, it is such a fine powder that it appears to dissolve in the water. It is used in the laundry to correct the yellowish tint derived from the ferrous bi- carbonate in the water (p. 400). If the goods are not freed by rinsing from soap and soda, however, the alkali liberated by hydrolysis of the latter enters into double decomposition with the Prussian blue and produces more rust: Fe 4 (Fe(CN) 6 ) 3 + 12NaOH-4Fe(OH) 3 | + 3Na 4 Fe(CN) 6 . IRON Fe 467 Potassium Ferricyanide K 3 l Fe(CN) & 111 . The difference be- tween this and the preceding salt can be seen by writing the formulae thus: 4KCN,Fe(CN) 2 and 3KCN,Fe(CN) 3 . In the ferricyanide the iron is trivalent and the negative ion Fe(CN) 6 = ~ is also trivalent as a whole. It is a soluble salt, of brown color, made by oxidizing the ferrocyanide. With ferric salts it gives only a brown solution, but with ferrous salts it gives a deep blue precipitate of ferrous ferricyanide Turnbulls blue : 2K3Fe(CN) 6 -*Fe 3 (Fe(CN) 6 ) 2 | + 6KC1. Blue-Prints. Some ferric salts are reduced to ferrous salts when exposed to light. Thus ferric oxalate Fe 2 (C 2 O4) 3 will keep in the dark, but in light gives ferrous oxalate FeC 2 O4i Fe 2 (C 2 O 4 ) 3 ->2FeC 2 O 4 + 2C0 2 | - When paper is dipped in ferric oxalate solution and dried, and a fern (or ink drawing on transparent paper) is placed over the prepared sheet, sunlight will reduce the iron to the ferrous condi- tion excepting where the fern or ink lines protect it from the light. When the sheet is now dipped in potassium ferricyanide solution (developer), the ferric oxalate gives only the brown substance which can be washed out. But the parts exposed to the light turn 1 deep blue from the precipitation of ferrous ferricyanide in the paper. The pattern is white on a blue ground. In regular blue- print paper ammonium-ferric citrate takes the place of the oxa- late, and the ferricyanide has already been applied to the paper, so that only exposure and washing remain to be done. Ink. Writing ink is commonly made by adding ferrous sulphate to an extract of nut-galls, giving ferrous tannate. This salt is soluble and almost colorless, but when spread on paper it is slowly oxidized to the insoluble, black ferric tannate. To make the writing visible from the first, a blue or black dye is added to the ink. 468 TEXT BOOK OF CHEMISTRY Ink, Rust 9 and Other Stains. Fresh writing ink, being soluble, can usually be washed out with water, if the latter is used at once. After the oxidation (p. 467) has occurred, the ferric tannate must be reduced again, by soaking the part for 12 hours or longer in ammonium oxalate solution, and the ink can then be washed out. Rust stains are often rendered soluble by ammonium oxalate also. Applying hydrochloric acid from a dropping tube, while the cloth is stretched loosely over a cup of hot water, dissolves the rust slowly as ferric chloride. The spot is dipped in the water, between applications of the acid, and finally rinsed with weak ammonia. Coffee and tea stains can be washed out by boiling water, instantly applied. Old stains of tea, coffee, or fruit, and stains from dyes, can be removed only by very dilute hypochlorous acid (mix solutions of bleaching powder and vinegar). Stains on silk can hardly ever be completely removed, and hypochlorous acid must never be applied to them. Sugar and syrup stains are removed by water. For grease, stretch the stained part on blotting paper (to absorb the solution and prevent spreading) and apply the solvent (p. 108) in a circle round the spot. For mixed stains use the grease solvent first, then, after it has evaporated, apply the water or aqueous solution. (For stains produced by indelible ink, see p. 490; by paint, see p. 476; by grease, see p. 108.) Other Salts of Iron. Ammonium sulphide solution added to a ferrous salt gives a black precipitate of ferrous sulphide FeS. Soluble sulphides blacken the fingers because of the iron, derived from decomposition of proteins, which is present in perspiration. We can distinguish ferrous-ion Fe+ + from ferric-ion Fe +++ by the reaction with potassium ferricyanide, which gives a precipitate (blue) with the former only. An equally sharp distinction is obtained by adding potassium sulphocyanate, for, although the COBALT CO AND NICKEL N! 469 ferrous and ferric sulphocyanates are both soluble, the latter is blood red in color. The borax bead, with a particle of an iron compound, gives a green color (ferrous borate) in the reducing part of the flame and a yellow color (ferric borate) in the oxidizing part. COBALT Co AND NICKEL Ni These metals, like iron, are attracted by a magnet. Cobalt, like iron, has two series of compounds, in which it is bivalent and trivalent, respectively. In its salts, nickel is bivalent only, but the oxide and hydroxide Ni 2 3 and Ni(OH) 3 are also known. Cobalt. Cobaltite CoAsS is one of the chief ores. It is roasted to oxidize the arsenic and sulphur and the oxide of cobalt remains. There are no applications of the metal. Salts are made by the action of acids on the oxide, which is basic. The metal has a silvery luster, tinged faintly with pink, and does not tarnish. It displaces hydrogen very slowly from dilute acids, but is acted upon rapidly by nitric acid. Salts of Cobalt. Cobaltous chloride CoCl 2 ,6H 2 O is red in color, and when partially or wholly dehydrated becomes deep blue. Writing made with a dilute solution of this salt leaves pink traces so faint as to be invisible. But, when the paper is warmed, the hexahydrate is decomposed, and the writing appears blue. When the breath is now blown on the writing, it disappears once more (sympathetic ink). Cobaltous nitrate Co(NO 3 ) 2 .6H 2 is also a familiar red-colored salt. Cobaltous. sulphide CoS is precipitated (black) by ammonium sulphide, and is not soluble in dilute hydro- chloric acid. Cobalt aluminate Co 3 (AlO 3 ) 2 (blue) and cobalt zincate CoZnO 2 (green) affords tests for aluminium and zinc. Borax (bead) gives a deep blue cobalt borate. Metallurgy of Nickel. Nickel occurs in all iron meteorites. The chief sources of nickel are garnierite (a silicate) from New 470 TEXT BOOK OF CHEMISTRY Caledonia and pentlandite (Ni, Cu, Fe)S from Ontario. In the latter case, the ore is roasted and smelted and finally bessemerized (p. 460) . The result is an alloy of nickel and copper which is much used for sheet metal work under the name of monel metal. Pure nickel is separated from the copper by an electrolytic process (see copper, p. 485), or by the Mond process (see nickel carbonyl, p. 470). The production of nickel was, in Canada $10,230,000 (1911) and $13,450,000 (1912), and in the U. S. $127,000 (1911). Properties and Uses of the Metal. Nickel is a white, hard, malleable metal which takes, and keeps, a high polish even in moist air. Nickel-plating, deposited electrolytically on iron, has the same qualities. The metal is used also in alloys such as monel metal (copper, nickel, approximately 1:1), German silver (copper, zinc, nickel, 2:1:1), and nickel coinage (copper, nickel, 3 : 1). In alloys, although the red color of the copper is completely lost, nickel and copper are simply dissolved and not combined. Zinc and copper, however, give a compound Cu 2 Zn 3 . Nickel steel is used for armor-plate. Compounds of Nickel. These salts are green, and give green solutions (Ni++). The sulphate NiSO 4 ,7H 2 is familiar, as is also the double salt, ammonium-nickel sulphate (NH 4 ) 2 SO 4 ,- NiS0 4 .6H 2 0, used in nickel-plating. The sulphide NiS (black) behaves like cobalt sulphide. Nickel carbonyl Ni(CO) 4 is a vola- tile, colorless liquid (b.-p. 43). It is formed by passing carbon monoxide over warm, finely divided nickel, and is decomposed again, yielding nickel 99.6 per cent pure, by heating at 180. These actions are the basis of the Mond process for separating nickel from the alloy obtained from pentlandite (p. 470). The borax bead is reddish-brown (nickel borate) when oxidized, and colorless, but turbid (metallic nickel), when reduced. Exercises. 1. Why does paint protect iron from rusting? 2. Why does iron, in time, turn completely into rust, while zinc and aluminium receive only a slight film of tarnish? COBALT CO AND NICKEL N! 471 3. How could it be ascertained that ferric hydroxide is in col- loidal suspension, and not dissolved? 4. Make equations for: (a) the interaction of potassium ferro- cyanide and cupric sulphate CuS0 4 , giving a brown precipitate of cupric ferrocyanide; (b) the action of hydrochloric acid on ferric hydroxide (or rust). 5. In removing a mixed stain, why allow carbon tetrachloride (e.g.) to evaporate, before applying water (to remove sugar, e.g.) ? CHAPTER XXXIX LEAD AND TIN LEAD and tin are the best-known metallic members of the family, to which the non-metals silicon and carbon also belong. In their compounds they are bivalent or quadrivalent. LEAD PB Metallurgy of Lead. The chief ore of lead is galenite PbS. The ore, if rich, is roasted in a reverberatory furnace (Fig. 107, p. 458) until a part has been converted into the oxide PbO and sul- phate PbSC>4. The air is then shut off, and the temperature raised, so that these products may oxidize the remaining galenite: PbS + 2PbO - 3Pb + S0 2 T PbS + PbS0 4 -> 2Pb + 2S0 2 T The melted lead flows out. When the ore is poor, it is treated at a red heat for 24 hours with lime and air. The product, a mixture of PbS0 4 (with the calcium sulphate unchanged) is then reduced with coke in a closed furnace. The lead contains silver and some gold, which are recovered by Parke's process (see silver, p. 488). Properties of Lead. Lead is a soft malleable metal, with a bluish-white or grey luster when freshly cut. It is quickly, but only superficially, oxidized by the air. It acts very slowly upon hydrochloric acid, but not upon sulphuric acid. With nitric acid it gives lead nitrate Pb(N0 3 )2 and oxides of nitrogen. Soft water, in presence of air, dissolves it in appreciable amounts as hydroxide Pb(OH) 2 , and carbonic acid assists the process. Hard water, 472 LEAD Pb 473 however, produces a skin of carbonate and sulphate (insoluble) which protects the surface. Hence, lead pipes may be used for hard water, but not for rain water. In presence of air, acids (even feeble vegetable acids) interact with the metal, which is therefore unsuitable for kitchen utensils. The compounds of lead are poisonous and, not being eliminated, accumulate in the system. For this reason the manufacture and use of white lead are forbidden in France (laws of 1906 and 1912), and are subject to detailed regulation in Germany. Uses of Lead. The metal is rolled into sheets, which are used for lining tanks. Lead pipes are made by pressing the metal, while hot, through dies. Their pliability, and the ease with which they can be cut and soldered, fits them for use in plumbing and for covering electric cables. Type metal, containing 20 per cent of antimony, is harder than lead, and expands on solidifying. Small shot (p. 340) contain 0.5 per cent of arsenic. Solder (lead, tin 1 : 1) remains melted at a lower temperature than pure lead (m.-p. 326) and so can be applied to a lead joint without danger of melting the lead itself. Pure melted lead would almost always be above 326, and would melt the leaden object to be soldered (p. 479, Ex. 2). Oxides. Lead monoxide PbO is made by oxidizing melted lead in a current of air. At low temperatures a buff colored powder, massicot, is obtained. When the oxide is allowed to melt, it solidifies to a reddish-yellow, scaley mass litharge. The oxide is predominantly basic, absorbing carbon dioxide from the air and, with acids, giving salts. It is used in making glass and enamels. Stone and glass can be cemented with a mixture of massicot and glycerine. Minium or red lead Pb 3 04 is a bright-red powder formed by oxidizing the monoxide in air at 470 to 480. 474 TEXT BOOK OF CHEMISTRY Red lead is used in making flint glass and as the basis of the red paint (see p. 475) applied to iron. Moistened with linseed oil (p. 475), it is used by plumbers to make joints tight. Lead dioxide Pb0 2 , is a brown powder, made by treating red lead with diluted nitric acid. Lead Acetate. Lead monoxide dissolves in boiling vinegar or acetic acid to give the acetate (sugar of lead), a white soluble substance : PbO + 2HC0 2 CH 3 ^Pb(CO 2 CH 3 ) 2 + H 2 O. Additional lead oxide can be dissolved in the neutral solution to give a basic acetate, 2Pb(C0 2 CH 3 ) 2 ,Pb(OH) 2 . White Lead. The ordinary carbonate PbCO 3 has no appli- cations. White lead is a basic carbonate 2PbCO 3 , Pb(OHj 2 . It is a heavy, white, insoluble, amorphous substance. Mixed with linseed oil, it forms a white paint valued for its "body " or covering power (opacity). Other pigments are added to give the required color. Its disadvantage is the darkening, due to formation of the black lead sulphide PbS, which is produced by the hydrogen sul- phide in the air. The old Dutch process is still used extensively. Gratings ("buckles") of lead are placed above a little vinegar in small pots. The pots are covered with boards, on which manure or spent tan bark is spread. Other tiers of pots, boards and bark are placed on the first, until the shed is full. Carbon dioxide, warmth and moisture are furnished by the decaying bark. The gratings, by the end of three months, are converted into cakes of white lead : 2Pb + 2C0 2 + O 2 -> 2PbC0 3 X 2 2Pb + 2H 2 O + O 2 -> 2Pb(OH) 2 6Pb + 4CO 2 + 2H 2 O + 3O 2 -+ 2(2PbCO 3 ,Pb(OH) 2 ). The vinegar may be regarded as a catalytic agent. It produces the basic acetate (p. 474), which is decomposed by the carbonic LEAD Pb 475 acid with liberation of the acetic acid once more. The white lead is ground, first under water, and then in oil. In Mild's process (used e.g. by the Acme Co., Detroit) melted lead is atomized by a jet of steam, and the lead dust is beaten with vinegar, air, and carbon dioxide in a vat for about seven days. In the French process white lead is precipitated by a stream of carbon dioxide from a solution of the basic acetate. Other Compounds of Lead. Lead chromate PbCrC^ is precipitated by adding potassium chromate (or dichromate) solu- tion to a solution of a salt of lead. It is used as a pigment (chrome- yellow). Lead chloride PbCl 2 (white) is very little soluble in cold water and the iodide PbI 2 (yellow) is insoluble. Both are formed by precipitation. Lead sulphide PbS (black) is precipitated by hydrogen sulphide, even from acid solutions (see picture restoring, p. 304). The sulphate PbSO 4 is the most insoluble salt. On this account, the workmen in white lead works add a little sulphuric acid to the water they drink. Zinc, or any of the metals more active than lead, when placed in a solution of a soluble salt of lead, will displace the metal, and deposit it in a mossy form (' 1? ad-tree"): Zn + Pb++-> Zn++ -f Pb j . Paints. A paint consists of a suspended solid powder, and a " drying" oil which will harden to a tough resin on exposure to the air and attach itself securely to the surface. The solid powder (pigment) gives opacity and, when desired, color to the paint. Since it does not shrink, it also "fills" the paint, and prevents the formation of minute pores which otherwise would be found in the " dried" oil. The pigment is thus necessary to enable the paint to protect the object from the action of the oxygen and moisture in the air. The " dryer" does not dry by evaporation, but gives a resin by oxidation. Linseed oil and hemp oil are examples. They contain 476 TEXT BOOK OF CHEMISTRY glyceryl esters of unsaturated acids, such as that of linoleic acid C 3 H5(CO 2 Ci7H3i)3 which contains four units of hydrogen less than stearic acid (p. 371). Boiling before use makes the esters more rapidly oxidizable by air. The cement of red lead and oil (p. 473) sets rapidly because the red lead itself oxidizes the oil. Removal of Paint Spots. Paint spots are removed by backing the material with blotting paper and applying turpentine (not to silk) or benzene to dissolve the resin. The loosened pig- ment can then be brushed out. Varnish, since it is made by dis- solving a resin in alcohol, can be removed with alcohol. Pitch (hydrocarbons) is removed by kerosene, because all hydrocarbons are mutually soluble. The Storage Battery. In the ordinary lead accumulator the plates consist of leaden gratings. The openings are filled with finely divided lead in one plate and with lead dioxide in the other. PbO, DISCHARGE 2H+ > -S0 4 = PbSO 4 CHARGE +-2HH 804= - FIG. 110. FIG. 111. These, and the dilute sulphuric acid in the cell, are the active sub- stances when the cell is charged. When the battery is used, the S0 4 = ions migrate towards the plates filled with the lead (Fig. 110), and convert this into a mass of the insoluble lead sulphate: LEAD Pb 477 SOr + Pb -> PbSO 4 + 26. These plates receive the negative charges. Simultaneously, the H+ ions move towards the other plates and there reduce to monoxide the lead dioxide with which they are filled. Pb0 2 + 2H+ -> H 2 + PbO + 20 . These plates acquire positive charges and, by interaction of the lead monoxide with the sulphuric acid, become filled, like the negative plates, with lead sulphate. During the discharge, much sulphuric acid is thus removed from the cell fluid, and the approach- ing exhaustion of the cells can thus be ascertained by measuring the specific gravity of the fluid. The E.M.F. of the current is a little over 2 volts. The charging is done by passing a current through the cell, in the opposite direction to the one which it yields (Fig. 111). The H + ions are attracted to the negative plate and an equivalent number of S0 4 = ions are formed, so that only lead remains : PbSO 4 + 2H+ + 20 -* Pb + 2H+ + S0 4 = . Simultaneously, the SO 4 = is attracted by the positive plate and, with the lead sulphate there present, forms lead persulphate: SOr + PbS0 4 + 20 - Pb(SO 4 ) 2 . The persulphate, being a salt of quadrivalent lead, is at once hydrolyzed and the filling of this plate is thus changed into lead dioxide: Pb(SO 4 ) 2 + 2H 2 O > Pb0 2 + 2H 2 SO 4 . Both plates are thus brought back to the con- dition in which they were before the discharge. The last set of charges consumes energy, while the first set liberates energy. Both may be stated in a single equation: charge > 2PbSO 4 + 2H 2 O <=> Pb + 2H 2 SO 4 + PbO 2 . < discharge In the Edison cell, when charged, one plate is of iron and the other contains nickelic oxide Ni 2 Oa. The cell liquid is a solution of potassium hydroxide. When the cell operates, the nickelic 478 TEXT BOOK OF CHEMISTRY oxide is reduced to Ni(OH) 2 and the iron is oxidized to Fe(OH) 2 , an action which delivers energy: Fe + 3H 2 + Ni 2 3 ^ Fe(OH) 2 + 2Ni(OH) 2 . When the cell is charged, the nickel is reoxidized and the iron reduced. TIN SN Metallurgy. Tin is obtained from cassiterite or tin-stone SnO 2 (stannic oxide). The largest supply comes from the East Indies, the next largest from Bolivia, and the third largest from Cornwall. The ore is roasted to expel sulphur and arsenic, and reduced with coal in the reverberatory furnace. The melted metal is cast in blocks ("block tin"). The metal was well known to the ancients (found in Egyptian tombs). Properties. The metal is white, and markedly crystalline. It is soft and malleable (tinfoil), and melts at 232. Tin does not tarnish in the air. With concentrated acids it acts rapidly. Hydrochloric acid gives stannous chloride SnCl2 and hydrogen. Concentrated sulphuric acid gives stannous sul- phate SnSCX, sulphur dioxide (p. 286) and water: Sn + 2H 2 S0 4 -> SnS0 4 + S0 2 t + 2H 2 0. Nitric acid gives the insoluble, white meta-stannic acid H 2 SnO 3 : 4HN0 3 + Sn - H 2 Sn0 3 i + 4N0 2 + H 2 O. Uses. Tin plate, used in making "tin" cans, is produced by dipping cleaned sheets of mild steel in melted tin. So long as the layer of tin remains intact, the iron is protected from rusting. But, if the tin layer is damaged, the iron rusts. The iron being the more active metal of the two, the tin acts as a contact agent and actually hastens the rusting. Scrap tin plate from can factories, bearing 3 to 5 per cent of tin, is "detinned." In the Goldschmidt process the scrap is cleaned, dried, and exposed to dry chlorine, which converts the tin into the TIN sn 479 liquid stannic chloride SnCl 4 , but leaves the iron unaffected. The chloride is used in mordanting. Copper vessels for cooking and brass pins are also coated with tin, to preserve them from the action of air and moisture. Tin pipes are used where lead would be unsafe, as, for example, for soda-water and beer. Tin enters into many useful alloys, such as bronze (with copper) and solder. Compounds of Tin. Stannous chloride SnCl2,2H 2 0, "tin crystals," and a hydrate of stannic chloride SnCl4,5H 2 are useo) in mordanting. With soda the former gives stannous hydroxide Sn(OH) 2 and the latter stannic acid Sn(OH) 4 , both gelatinous substances which are precipitated in the goods to be dyed. Stan- nic acid is also sometimes precipitated in flannelette (a material made of cotton) to render it non-inflammable, and always in silk to "weight" it (the increase may be from 25 per cent to 300 per cent or more). Stannous sulphide SnS (brown) and stannic sulphide SnS2 (yellow) are precipitated by hydrogen sulphide, even from acid solutions of stannous and stannic salts, respectively. They both interact with a solution of ammonium polysulphide and dissolve. Exercises. 1. Make equations for: (a) the action of air and water on lead to form the hydroxide; (b) the formation of basic lead acetate by dissolving lead oxide in lead acetate solution; (c) the action of hydrogen sulphide on white lead; (d) the precipita- tion of PbCl 2 and of PbI 2 . 2. Tin melts at 232 and lead at 326. Solder (1 : 1) melts (and solidifies) at 210. Why is this? What is the advantage of using solder? 3. Make equation for: (a) the action of soda on solutions of stannous and stannic chlorides, respectively: (b). the action of hydrogen sulphide on the same salts. CHAPTER XL COPPER AND MERCURY IN this chapter we encounter the first metals (if we except As and Sb) which are below hydrogen in the activity list, and do not displace that element from dilute acids. Copper and mercury both have two valences, so that we have cupric-ion Cu"^ 1 " and cup- rous-ion Cu + and mercuric-ion Hg++ and mercurous-ion Hg+. All the soluble compounds of both are poisonous. COPPER Cu Occurrence. Copper occurs free in considerable amounts, particularly on the Michigan shore of Lake Superior. Cuprous oxide Cu2O and basic carbonates, like malachite CuC0 3 ,Cu(OH) 2 , are less common. The latter is often used as an ornamental stone. A large proportion of commercial copper is obtained from chal- copyrite Cu 2 S.Fe 2 S 3 . Metallurgy. The free copper, after being concentrated (freed from gangue) by washing, is smelted with a flux. The car- bonate is roasted, leaving the oxide. The oxides are reduced with coal. The sulphide ores are more difficult to reduce, and the presence of so much iron complicates the process. They are first roasted. This removes much of the sulphur as sulphur dioxide, leaving Cu 2 and Fe20 3 . Next the roasted material is treated in a blast furnace, along with " green" (unroasted) ore, sand (if silica is not present in the ore), and coke. Some of the iron is oxidized and removed in the. slag (as silicate). The product, known as copper, "matte," is a mixture of cuprous sulphide Cu 2 S with ferrous sul- 480 COPPER CU 481 phide FeS. The third stage is to bessemerize the melted matte with sand in a converter. Here the rest of the iron is oxidized and eliminated as silicate in the slag, and the sulphur escapes as SC>2. The slag and metallic copper are poured separately. The latter gives off some dissolved sulphur dioxide in bubbles as it solidifies and, from its appearance, is named blister copper. Finally, since the copper now contains dissolved cuprous oxide Cu 2 O, the blister copper is melted and "poled," by stirring with green wood. The gases (hydrocarbons, etc., p. 435) given off by the heated wood reduce the oxide to copper. If the copper is to be refined electro- lytically (p. 485), it is then cast in plates 3 feet square and f inch thick. Properties. Copper has a characteristic bright yellow-pink luster, quickly darkened by oxidation. It is second only to silver in electrical conductivity and to iron in tenacity. It is third in order of malleability ("gold-leaf"). It melts at 1057. In moist air copper acquires a green coating of basic carbonate, which protects the metal. It is little affected by dilute hydro- chloric or sulphuric adds, when air is excluded. Hot concentrated sulphuric acid gives cupric sulphate and sulphur dioxide (p. 275), and nitric acid gives cupric nitrate and oxides of nitrogen (p. 317). Uses. Pure copper is used for electric wires and cables. Traces of other metals greatly reduce the conductivity. Kettles, stills, and evaporating pans are made of copper. It is used for sheathing ships and for bolts, because it resists corrosion by sea water. It enters into important alloys, such as brass (18 to 40 per cent of zinc), and bronze for coins (4 per cent tin and 1 per cent zinc), for gunmetal (10 per cent tin), and for bell-metal (20 to 24 per cent tin). Aluminium bronze (5 to 10 per cent aluminium) is used for the hulls of yachts. All of these are composed in part of com- pounds, such as CuaSn and Cu 2 Zn 3 . 482 TEXT BOOK OF CHEMISTflY Cupric Sulphate CuSO 4 .5H 2 O. The hydrated sulphate, bluestone or blue vitriol is made by continuously agitating granu- lated copper with air and warm dilute sulphuric acid: 2Cu + 2 + 2H 2 S0 4 -> 2CuS0 4 + 2H 2 O. The Uue crystals form on strips of lead suspended in the warm, saturated solution. Because of slight hydrolysis, giving an active acid and a weak base Cu(OH) 2 , the aqueous solution is acid in reaction. Cupric sulphate is used in battery solutions. The salt is em- ployed in minute proportions to destroy algae, which otherwise confer a disagreeable taste and odor on water that has been stand- ing in reservoirs. Seed for cereals is moistened with a dilute solu- tion to prevent "smuts." It is used in making Bordeaux mixture (p. 482). Cupric Hydroxide Cu(OH) 2 . The hydroxide is a Hue, gelatinous precipitate, formed when an alkali is added to cupric sulphate solution. It is used as a mordant. A mixture of cupric sulphate solution and milk of lime (Ca(OH) 2 ), Bordeaux mix- ture containing this precipitate, is used extensively as a spray on grape vines and other plants, to prevent the growth of fungi. Cupric hydroxide dissolves in ammonium hydroxide solution, giving a liquid of deep blue color (containing Cu(NH 3 ) 4 (OH) 2 ) which is employed as a solvent for cellulose. It dissolves also in a solution of Rochelle salt (p. 221), giving Fehling's solution, a reagent used in testing for glucose and similar reducing agents. When this reagent is added to a liquid containing glucose (p. 221), red cuprous oxide Cu 2 O is precipitated. Cuprous Oxide Cu 2 O. This oxide, mixed with CuO, is formed by gentle heating of copper in air, and is best prepared by use of Fehling's solution (p. 482). It is employed in making ruby glass and in coloring porcelain. COPPER cu 483 Cupric Oxide CuO. When the liquid containing the blue precipitate of cupric hydroxide is boiled, the blue color changes to black and cupric oxide is thrown down: Cu(OH) 2 -+CuO + H 2 0. The same oxide is made by heating cupric nitrate (p. 318) or car- bonate. Organic Analysis. The oxide is used in the laboratory to ascertain the compositions (and formulae) of organic compounds (determination of carbon and hydrogen). A weighed amount of the organic compound is placed in a horizontal tube, between heated masses of the oxide. A stream of oxygen or air carries the vapor of the organic compound over the cupric oxide, which oxidizes it to water and carbon dioxide. The first is absorbed in a weighed U-tube filled with calcium chloride, and the second is caught in a weighed vessel containing potassium hydroxide. From the increase in weight in each case, the corresponding weights of hydrogen and carbon (derived from the weighed portion of the organic compound) are calculated. Other Compounds of Copper. Hydrogen sulphide precip- itates cupric sulphide CuS (brownish-black) even from acid solutions of cupric salts. When a particle of a compound of copper is fused with borax (bead), blue cupric borate is formed. More active metals, such as zinc or iron, displace copper from solutions of its salts, so that a blade of a knife, for example, receives instantly a red coating of copper when immersed in such a solution : Fe + Cu++->Fe++ + Cu | . Copper Plating. When platinum or carbon plates, connected with a battery, are immersed in a solution of cupric sulphate, copper is deposited on the negative plate (cathode). The S0 4 = migrates towards the positive wire (anode) and there produces oxygen and sulphuric acid: 2SO 4 + 2H 2 -> 2H 2 S0 4 + O 2 T 484 TEXT BOOK OF CHEMISTRY If the anode is made of copper itself, however, the S04 = migrates, but is not discharged. Instead, copper goes into solution (Fig. 112) as Cu++, in amount equal to that deposited on the other plate. Thus the quantity of cupric sulphate in solution remains unchanged, and the effect is, virtually, to transfer copper from the copper anode to the cathode. Electrotypes. A copper electro- type of an object like a medal is made by first preparing a cast of the medal in plaster of Paris or wax. The sur- face of this is rubbed with graphite, to render it a conductor, and the cast is used as the cathode in a cell like that just de- scribed. The deposit of copper, when stripped off, is found to show an exact reproduction of the engraving, etc., on the object. Book plates are made by taking a cast of each page of type, preparing the copper electrotype, and then strengthening and thickening it by filling the back with melted lead. The printing is then done from the electrotype. For newspapers this process is too slow, and the plate is made from the cast by means of melted stereotype-metal (lead, antimony, tin; 82 : 15 : 3). Copper Refining. The copper, as obtained from the ores by the treatment already described (p. 481), contains a certain amount of silver, gold, and baser metals. The former pay for the cost of refining, and the simultaneous removal of the latter give pure copper suitable for electrical purposes. The principle is the same as that used in electroplating. The heavy plates of poled copper (p. 481) are hung at intervals in large, lead-lined vats of copper sulphate solution and form the anodes (Fig. 113, diagrammatic, view from above). The metal is de- posited on thin sheets of copper, which are coated with graphite to permit the deposit to be easily stripped off. These sheets hang MERCURY Hg 485 in the vat between the anodes, and are connected with the negative wire. The copper, along with such traces of more active metals, like zinc (see activity list), as are present, is ionized and goes into solution, until the anode is reduced to a + skeleton and is exchanged for a fresh one. The less active metals, such as silver and gold, as well as traces of sulphides, are not ionized. They fall to the bottom of the vat, as a sort of heavy mud. At the cathode copper ions alone are discharged, and deposited, because copper is the least active of the metals present in ionic form. In this way copper, 99.8 per cent pure, can be obtained, and gold and silver are recovered from the mud. FIG. 113. Nickel Plating. Here the bath contains an ammoniacal solution of ammonium-nickel sulphate (p. 470), and a plate of nickel forms the anode. The article to be plated is carefully cleaned, to secure a uniform deposit, and is suspended as cathode in the vat. The surface of the deposit is afterwards burnished. MERCURY HG Metallurgy. Mercury occurs both free and as mercuric sul- phide HgS, cinnabar, and usually in the same localities. Most of the ore comes from California and Spain (Tarshish). Decoration with "rubies of Tarshish" refers to tattooing with powdered cin- nabar (red). The ore is roasted, sulphur dioxide escapes, and the vapor of mercury is condensed in long, tortuous flues. Properties. Mercury is a liquid at ordinary temperatures, hence its name, quicksilver (i.e. live silver). It freezes at 40, and boils, giving an invisible, non-conducting vapor, at 357, The 486 TEXT BOOK OF CHEMISTRY vapor density shows the molecules to be monatomic, as indeed are the molecules of all metals of which the densities have been measured. The metal has a silvery metallic luster, which is not affected by the air, and a high specific gravity. Mercury dissolves other metals, forming alloys or amalgams. Mercury vapor, and mercury when finely divided, are poisonous. Small amounts are triturated with honey and mixed with pulverized roots in making blue pills, and with greasy substances in making mercurial ointments. Mercury, when moderately heated, combines with oxygen, form- ing mercuric oxide (red), but the action is reversible, and the oxide is decomposed by stronger heating (p. 15). It combines readily with sulphur and the halogens. Dilute hydrochloric and sulphuric acids are not attacked by mercury. Concentrated nitric acid attacks and dissolves it readily. Uses. Mercury is used in filling thermometers and barom- eters. Sodium amalgam is used in the laboratory and the zinc plates of batteries are amalgamated superficially to protect them when the battery is not in use. Dentists fill teeth with mixtures of mercury with silver, copper, cadmium and other metals, which quickly set to a solid amalgam. The pulverized ores of gold and silver, mixed with water, are allowed to trickle over thin layers of mercury. The latter dissolve the particles of the precious metals, while the sand passes on. Compounds of Mercury. As in the case of copper, there are two sets of compounds mercurous Hg 1 and mercuric Hg n . All the common salts are completely volatile, with or without decom- position, when strongly heated (vapor poisonous!). Mercuric chloride or corrosive sublimate HgCl2 is made by subliming a mixture of mercuric sulphate and sodium chloride. It is a white substance, soluble in water. Hydrogen sulphide MERCURY Hg 487 precipitates mercuric sulphide HgS (black) from the solution. Mercuric chloride is a violent poison. A very dilute solution is used as an antiseptic. Great precaution should be taken, however, to make impossible the accidental swallowing of the antiseptic tabloids, as each contains enough of the material for several fatal doses. Albumen (white of eggs) forms insoluble compounds with it, and is used as an antidote. Mercurous chloride or calomel HgCl is precipitated as an amorphous white powder when a chloride is added to a solution of a mercurous salt. It is used in medicine to stimulate all organs producing secretions. Ammonium hydroxide changes calomel to a black mixture of free mercury and mercuric amidochloride Hg(NH 2 )Cl. Mercuric fulminate Hg(ONC)2 is used in percussion caps (p. 383). Baser metals precipitate mercury from solutions of its salts. The grey deposit is best seen on a clean strip of copper foil: Cu + Hg++ -+ Cu++ + Hg | Exercises. 1. Make an equation: (a) for the oxidation of ethyl alcohol by heated cupric oxide; (b) for the precipitation of cupric sulphide from cupric sulphate solution. 2. When we electrolyzed potassium chloride solution (p. 484), hydrogen was liberated at the cathode. What principle, used in the electrolytic refining of copper, does this phenomenon illustrate ? 3. How could you recognize cupric sulphate solution by showing that it contained, (a) cupric-ion, (b) sulphate-ion (p. 286) ? 4. Make an equation for the liberation of mercury from cinna- bar. 5. What gases should you collect over mercury? 6. What other salts, beside those of mercury, are volatile ? 7. How should you recognize a salt of mercury in solution? CHAPTER XLI * SILVER, GOLD, PLATINUM DESCENDING the activity list, we now reach the "noble" metals, which are the least active. They do not combine with oxygen, even when heated. Silver and gold have long been in use. SILVER AG. Occurrence. Native silver is found in many localities, usually in small amounts. The chief supply of the metal is obtained from the ores of lead and copper, which contain silver sulphide Ag2S. The chief localities are in California, Australia, and Mexico. Metallurgy. After the lead (bearing silver) has been extracted from the ore (p. 472), it is melted in large caldrons, a small propor- tion of zinc is added, and the whole is vigorously stirred (Parke's process). Zinc is only slightly soluble in lead, but it combines with silver in several proportions. The zinc-silver alloy rises to the surface, solidifies (while the lead is still molten), partly as alloy and partly as a compound (usually Ag 2 Zn 5 ), and is skimmed off. The most of the adhering lead is pressed out, and the compound (or mixture) is placed in graphite retorts, in which the zinc is re- moved by distillation. The silver and lead which remain are heated in a blast of air (cupelled) to oxidize the lead. The melted litharge flows off and the silver is then cast. The gold, which accompanies the silver through this treat- ment, is separated electrolytically (see copper, p. 484). The silver-gold alloy forms the anode, and silver nitrate the vat-liquid. The silver, being the more active metal, is ionized and deposited on the cathode, while the gold collects as a powder in a bag which surrounds the cathode, 488 SILVER Ag 489 An older method, still used, is to heat the silver with concen- trated sulphuric or nitric acid, either of which will dissolve the silver and leave the gold. From the solution the silver is dis- placed by the action of plates of copper: Cu + 2Ag+ - Cu++ + 2Ag | Over sixty million ounces (troy) of silver are obtained annually in the United States (1911), this being about one-third of the world's production. Properties. Silver is fairly hard, considering its great ductility and malleability. It is the best conductor of electricity. When an electric discharge passes between the ends of two silver wires, held under water, silver is dispersed at the points and forms a colloidal solution. The color of the solution varies from brownish to pink, according to the conditions. Colloidal solutions of gold and platinum can be made in the same way. The metal is oxidized by ozone (p. 302), although not by oxygen. Sulphur compounds in the air tarnish the surface (Ag2S, for details, see p. 270), as do also eggs, secretions from the skin (see proteins), and rubber which contains sulphur. Cold nitric acid and hot concentrated sulphuric acid are at- tacked by it, giving the nitrate and oxides of nitrogen, and the sulphate and sulphur dioxide, respectively: 3Ag + 4HNO 3 (da.) -> 3AgNO 3 + 2H 2 + NO t Uses. For silver ware and coins the metal is alloyed with copper. American and Continental coins contain 90 per cent of silver ("900 fine")- British coins are 925 fine, which is the proportion in " sterling silver." Articles of baser metal are plated with silver. The bath contains potassium argenticyanide KAg(CN) 2 , made by adding potassium cyanide in excess to silver nitrate solution: KNC + AgN0 3 - KN0 3 + AgCN. 490 TEXT BOOK OF CHEMISTRY This solution gives a compact deposit. The anode is of silver, so that the silver in the solution is replenished as quickly as it is de- posited. Mirrors are silvered by cleaning the surface and pouring over it a solution containing silver nitrate, ammonium hydroxide, and a reducing agent like formaldehyde CH 2 0, or grape sugar : 4AgOH -j- CH 2 - 3H 2 + 4Ag | + C0 2 . The film of silver adheres to the glass and is washed, dried, and varnished. Silver Nitrate AgNO 3 . This salt is deposited from solution (p. 489) in colorless crystals. Its solution is neutral in reaction, which shows that silver hydroxide is not a weak base. It melts easily and is cast in sticks (lunar caustic), which are used in cauterizing sores. It is the chief source of the other compounds of silver. It is used in some hair dyes, and in indelible ink. In the latter case the organic matter in the goods reduces it, with the help of light, to metallic silver. Sodium thiosulphate solution (See p. 491) will very slowly remove the stain. Silver Halides. Silver chloride AgCl is precipitated (white) when a soluble chloride is added to a solution of a salt of silver: AgNO 3 + KC1 - AgCl | + KN0 3 . Silver bromide AgBr is precipitated with bromide-ion and silver iodide Agl with iodide-ion. These compounds have a yellowish tinge. The chloride and bromide are easily dissolved by am- monium hydroxide solution, giving the ion Ag(NH 3 ) 2 + , and also by sodium thiosulphate solution Na 2 S 2 3 ("hypo"). All the halides of silver are decomposed by light, which liberates the halogen, and finally leaves metallic silver. Photography. The taking of a photograph involves four processes preparation of the plate, exposure, development, and fixing. SILVER Ag 491 Silver bromide is precipitated in an aqueous suspension of gelatine. The mixture is kept warm, to permit the precipitate to become more sensitive to light by acquiring a coarser grain (" ripen"). The " emulsion" is applied to plates of glass or strips of transparent celluloid (films). The brief exposure to the image of the object, well-focussed in the camera, produces no visible effect. But the bromide is thereafter more easily reduced to metallic silver, in proportion to the intensity of the light that fell upon each part. Development consists in applying a reducing agent, of such slight activity that its effect during the process on non-illumi- nated parts of the bromide is practically zero. Ferrous oxalate, or an alkaline solution of pyrogaUic acid or of hydroquinone may be used. The reduction goes fastest and deposits most silver when the illumination was most intense. Thus, the plate be- comes most opaque where the object was brightest, and vice versa. On account of this reversal, the plate is called a negative. With the potassium salt of hydroquinone C6H 4 (OH) 2 , quinone CeH^ is formed: 2AgBr + C6H4(OK) 2 -* 2Ag + 2KBr + CeKA. The foregoing processes are all carried out in a faint red light, which is almost without action on silver bromide. To prevent the gradual reduction of the remaining, unchanged bromide to silver by daylight, it is dissolved out by soaking the plate in sodium thiosulphate (fixing) . The plate is now clear where no silver was deposited. The negative is finally washed thoroughly to re- move all except the gelatine and the silver image, and is then dried. In printing the prepared paper is illuminated through the nega- tive, and light and dark are again reversed. The denser parts of the negative protect the paper below them, and leave these parts white. On printing out paper, silver chloride suspended in egg albumen is the sensitive substance, and the silver is liberated in a reddish, colloidal condition. The color is improved by toning 492 TEXT BOOK OF CHEMISTRY with a solution containing gold chloride, as part of the silver goes into solution and gold (purplish) is deposited in its place. The print is fixed with hypo, washed, and dried. Developing papers like velox (invented by Baekeland) are essentially like plates (silver bromide in gelatine), and are exposed, developed, and fixed in the same way. The making of blue prints has already been described (p. 467). GOLD Au. Occurrence and Extraction. Gold is found in the free condition in veins of quartz in almost every part of the world. Alluvial deposits, resulting from the breaking up of such rock, contain gold also, and nuggets weighing many pounds have been discovered. It is found also in combination with tellurium, an element resembling sulphur. In vein mining (e.g. in the Transvaal) the rock is pulverized with iron stamps working in an iron trough. The powder is washed in the form of mud over plates of copper amalgamated with mercury, in which 55 per cent of the gold dissolves. The amalgam is afterwards scraped off, the mercury removed by dis- tillation, and the gold residue refined. The tailings still contain 45 per cent of the gold, adhering to the particles of rock. They are covered with sodium cyanide (p. 405) solution, and exposed to the air, until the gold has been dissolved as sodium aurocyanide NaAu(CN) 2 . From this solution the gold is deposited by electrolysis, or displaced by zinc. The alluvial deposits are washed, on a small scale, in "cradles" (shallow pans) and, on a large scale, by being carried by water down a long trough (Placer Mining). The gold, having a much higher specific gravity than the rock, sinks to the bottom, while the rock particles are carried away. In the trough the gold settles between cleets nailed across the bottom. In hydraulic mining, a modification of placer mining, very heavy streams of water are thrown against the deposit. PLATINUM Pt 493 Over one-fourth of the world's supply comes from the United States (4,700,000 ounces in 1911) and a little less than one-fourth each from South Africa and Australia. Properties. Gold is yellow in color, copper being the only other familiar metal which is colored. It is the most malleable and ductile of metals. It melts at 1075. To enable it to resist wear, it is alloyed with copper. Pure gold is "24 carat" fine. British gold coins are 22 carat, and American, French and German coins 21.6 carat (90 per cent gold). Gold is not affected by oxygen or by hydrogen sulphide. It does not interact with any single acid. It combines directly, however, with chlorine and bromine. It dissolves in aqua regia (hydrochloric and nitric acids, mixed). This happens, not be- cause aqua regia is more active as an oxidizing agent than the sub- stances it contains, but because it oxidizes, and also furnishes the chloride-ion Cl~ required to produce the exceedingly stable nega- tive ion of chlorauric acid HAuCU, namely AuClT. Uses. Most of the metal is used in coins and bars as a medium of exchange. It is beaten into gold leaf. It is employed in making potassium chloraurate KAuCli for toning photographs. Gold plating on silver and other metals is carried out by using a gold anode and a bath of sodium aurocyanide NaAu(CN) 2 solution. PLATINUM Pt. Occurrence and Extraction. Platinum is found in the free condition in alluvial sand, and chiefly in the streams of the Ural Mountains and the Caucasus. The separation from osmium, indium, and other metals, which accompany it, is a complex operation. Properties. Platinum is a malleable and ductile metal with a greyish-white luster. It melts in the oxy-hydrogen flame (m.-p. about 1780), but not in the Bunsen flame. 494 TEXT BOOK OF CHEMISTRY The metal is not affected by air, -water, or acids, excepting aqua regia. In the latter instance, chloroplatinic acid H^PtCle is formed. Platinum absorbs hydrogen and oxygen. When finely divided, so as to present a large surface, it catalyzes powerfully many chemical actions (pp. 56, 60). The forms used are platinum black, a powder precipitated from chloroplatinic acid by an active metal (e.g. zinc) ; platinum sponge, a porous mass made by decom- posing ammonium chloroplatinate by heat ; and platinized asbestos, made by dipping the asbestos in chloroplatinic acid solution and heating. It unites, slowly, with chlorine and bromine. It combines also with carbon, phosphorus, and silicon, and alloys itself with many metals, so that reactions which liberate any of these elements can not be carried out in vessels of platinum. It also acts upon fused caustic alkalies giving platinates. Uses. Since the metal does not melt in the Bunsen flame, and is not affected by many substances, it is employed in laboratory operations in the forms of wire, foil, and crucibles. Much of the metal is used by dentists, and in photography. Having the same expansibility as glass, it is fused through the bottoms of incan- descent lamps, to connect the filament with the exterior. Since the metal is not oxidized by the air, even when heated, it is used for electrical contacts. The metal is employed also for jewelry and, in Russia, for coinage. An alloy with iridium, offering special resistance to both chemical and mechanical wear, is used for the international standards of length and weight. On account of the increasing demand, and the lack of any substitute, platinum is now worth $105.00 (1913) per ounce troy (31.02 g.), or five times the value of gold ($20.67 per ounce). Exercises. 1. Make an equation for: (a) the action of con- centrated sulphuric acid on silver; (b) the decomposition of silver chloride by light. PLATINUM Pt 495 2. Make an equation for: (a) the displacement of gold from potassium chloraurate by silver; (b) the decomposition of chloro- platinic acid, and of (c) ammonium chloroplatinate by heat; (d) the action of zinc on chloroplatinic acid. 3. What would be the advantages and disadvantages of using gold instead of platinum for crucibles ? CHAPTER XLII MANGANESE, CHROMIUM, AND RADIUM THE first metallic elements we considered form simple positive ions (e.g. Na + , Ca" 1 " 4 ") only. The last two, on the other hand, appear almost exclusively in complex, negative ions, just as do the non-metallic elements (AuCLT, Au(CN) 2 ~, PtCl 6 = ). Some of the intermediate metals can give negative ions (e.g. Zn0 2 = , A103 = ~), although in most of their compounds they are positive (Zn" 1 " 1 ", Al" 1 " 1 "*) . The first two in the present chapter are equally familiar in both roles. MANGANESE MN Ore and Preparation. The commonest ore of manganese is pyrolusite Mn02, a soft, black mineral. The metal is obtained in pure form by mixing the pulverized dioxide with aluminium (Goldschmidt's process, p. 424) in a clay crucible and starting the reaction with magnesium : 3Mn0 2 + 2A1 -> 3Mn + 2A1 2 3 . At the high temperature (over 3000) the molten manganese sinks to the bottom, and the alumina floats above it. Properties and Uses. Manganese is hard and crystalline, with a greyish-white luster. It is tarnished superficially by moist air. In fine powder it slowly displaces hydrogen from boiling water. It acts vigorously on dilute hydrochloric and sulphuric adds, giving manganous salts: Mn + 2HC1 - MnCl 2 + H 2 f . The preparation and uses of spiegeleisen and ferro-manganese (p. 459), and of manganese steel (p. 462) have already been de- 496 MANGANESE Mn 497 scribed. Traces of manganese seem to be essential to the growth of plants. Manganous and Manganic Salts. There is a complete series of manganous salts (MnCl 2 , MnSO 4 , MnCOs, etc.), in which the element forms the positive radical and is bivalent. They are all more or less distinctly pink in color. With the borax bead all manganese compounds give an am- ethyst (manganic borate, Mn m ) in the oxidizing flame, and a colorless bead (manganous borate) in the reducing flame. Manganese Dioxide MnO 2 . This oxide, composing the soft, black mineral pyrolusite, is the commercial source of the metal. It is an oxidizing agent, and is mixed with black paints to make them "dry" rapidly. For the same reason, and because it is a conductor, it is used as a depolarizer in dry batteries and in the Leclanche cell. It is added to molten glass to remove the green color due to compounds of iron (p. 348). Potassium Manganate K 2 MnO 4 . When manganese dioxide is fused in a porcelain crucible (or on a platinum wire) with potas- sium hydroxide and potassium nitrate (to furnish oxygen), a green mass of potassium manganate is obtained: 2KOH + Mn0 2 + (0) - K 2 Mn0 4 + H 2 0. This salt is very soluble in water, and can be secured in green crystals only by enclosing the vessel containing the solution along with another containing a drying agent (CaCl 2 or H 2 S0 4 ). In this salt (K 2 0,MnO 3 ) the manganese is sexivalent, and the man- ganate radical Mn0 4 = is bivalent. Potassium Permanganate MKnO 4 . If the green solution of the manganate be diluted with water, it changes to a purple- red color, manganese dioxide is precipitated, and the clear liquid gives, on evaporation, crystals of potassium permanganate: 3K 2 Mn0 4 + 2H 2 - 4KOH + 2KMnO 4 + MnO 2 | . 498 TEXT BOOK OF CHEMISTRY This salt crystallizes in needles which have a deep purple color with a greenish luster by reflected light. In this salt (K 2 0,Mn 2 07), the manganese is heptavalent, and the radical MnO 4 ~ is univalent. The purple color of the solution is the color of the ion Mn0 4 ~, while the ion MnO/" of the manganates is green. The solution of potassium permanganate, especially when, by addition of an acid, permanganic acid HMn0 4 has been formed, is an active oxidizing agent. We employed it in oxidizing hydro- chloric acid to make chlorine (p. 135). For the same reason a dilute solution of potassium permanganate is used as an antiseptic and disinfectant. CHROMIUM Cr. Ore and Preparation. The chief ore of chromium is chromite FeO, Cr 2 3 , imported from Asiatic Turkey and New Caledonia. The metal is easily obtained, by reducing chromic oxide Cr 2 3 with aluminium (Goldschmidt process, p. 424). Properties. Chromium is a lustrous crystalline metal. It does not tarnish in air. It displaces hydrogen, however, from dilute acids, giving chromous salts CrCl 2 , CrSO 4 , etc. Chromium (3 per cent) gives a hard steel, and is often used with nickel (p. 462) in armor plate. The strange alloys which, although composed entirely of active metals, are hardly affected even by boiling acids (including nitric acid), usually contain chromium (e.g. 60 per cent Cr. 36 per cent Fe, 4 per cent Mo, no C). Potassium Chromate K 2 CrO4. Powdered chromite is roasted with potash and lime: 4FeO,Cr 2 3 + 8K 2 C0 3 + 70 2 - 2Fe 2 3 + 8K 2 Cr0 4 + 8C0 2 t and the potassium chromate dissolved out of the residue. It is a yellow, soluble salt (the ion Cr0 4 = is yellow), with which in- soluble chromates of other metals, such as lead chromate PbCrO* can be precipitated. CHROMIUM cr 499 When the solution is acidified, it turns red, and potassium dichromate H 2 Cr 2 7 crystallizes on evaporation: 2K 2 CrO 4 + H 2 SO 4 - K 2 S0 4 + 2KHCrO 4 -* K 2 Cr 2 7 + H 2 O. Potassium Dichromate K 2 Cr 2 O 7 . This salt (see above), forms large orange-red crystals. It is related to the chromate (K 2 0,CrO 3 ) by containing a double proportion of the anhydride -K 2 O,2Cr0 3 . Like the chromate, the dichromate gives, by precipitation, the insoluble chromates. When potassium hydroxide is added slowly to potassium dichro- mate solution, until the red color has changed to yellow, potassium chromate is formed: K 2 Cr 2 O 7 + 2KOH -> 2K 2 CrO 4 + H 2 0. Potassium (or sodium) dichromate solution is employed, in- stead of tan-bark, in tanning kid and glove leathers. Its use reduces the tune required from 8 to 10 months to a few hours. A reducing agent is employed to precipitate chromic hydroxide Cr(OH) 3 (p. 500) in the leather. By addition of sulphuric acid to a concentrated solution of a dichromate, chromic anhydride CrO 3 , is precipitated in red crys- tals. Classes of Chromium Compounds. It will be noted that three classes of compounds have been illustrated: Chromous salts CrCl 2 , CrSO 4 . Cr bivalent. Chromic compounds Cr(OH) 3 and salts. Cr trivalent. Chromates K 2 Cr0 4 , K 2 Cr 2 O 7 (and CrO 3 ). Cr sexivalent. Chromic Sulphate Cr 2 (SO 4 ) 3 : Chrome- Alum. The chromic salts are made by reduction from chromates or dichro- mates. If potassium dichromate solution be acidified with sul- phuric acid, only dichromic acid H 2 Cr 2 O 7 is formed. But if a reducing agent, like sulphurous acid, is added, then chromic sul- phate is formed: 500 TEXT BOOK OF CHEMISTRY Partial (1) : K 2 Cr 2 7 + H 2 S0 4 -> H 2 Cr 2 7 + K 2 S0 4 . Partial (2): H 2 Cr 2 7 + 3H 2 S0 4 - Cr(SO 4 ) 3 + 4H 2 + (30). Partial (3) : (30) + 3H 2 SO 3 -> 3H 2 S0 4 . K 2 Cr 2 7 + H 2 S0 4 + 3H 2 S0 3 - Cr 2 (S0 4 ) 3 + K 2 SO 4 + 4H 2 O. Adding the three partial equations, and cancelling 3H 2 S0 4 from each side, we get the final equation. When the solution stands, purple octahedral crystals of chrome-alum K 2 S0 4 ,Cr 2 (S0 4 ) 3 ,24H 2 are formed (see Fig. 39, p. 111). In photography this salt is often added to the fixing bath. Chromic Hydroxide Cr(OH) 3 and Oxide Cr 2 O 3 . When an alkali is added to chrome-alum solution, chromic hydroxide Cr(OH) 3 (green) is precipitated: Cr 2 (S0 4 ) 3 + GNaOH -> 2Cr(OH) 3 | + 3Na 2 S0 4 . Chromic hydroxide is the substance precipitated in the leather (p. 499). It is used also as a mordant in calico printing. When the hydroxide is heated, chromic oxide Cr 2 3 remains as a green powder. Chromium compounds give with borax a green bead, containing chromic borate. RADIUM Ra. The Discovery of Radium. In 1896 Henri Becquerel dis- covered that a crystal of a salt of uranium could, in the dark, reduce the silver bromide on a photo- graphic plate, even when a sheet of black paper (impervious to light) was placed be- tween. Evidently a radiation, different from light, was given out by the salt. Next he discovered that an electrometer (Fig. 114), in which the gold leaves had been caused to separate by charging with x' IG. 114 electricity, lost its charge rapidly when a salt of uranium was brought near to the knob connected with the leaves. Evidently the salt rendered the air a conductor, and this FOG-TRACKS FROM RADIUM (C. T. R. WILSON) 1, 1 Paths of helium atoms. 3. Part of 2, enlarged. 4. Paths of electrons. RADIUM RE 501 permitted the escape of the electricity. These discoveries, in the hands of a multitude of observers, have led to the development of an entirely new branch of our science, namely radio-chemistry. Pierre Curie and Mme. Curie (ne. Marie Sklodovska) found that the mineral pitch blende, impure uranium oxide UaOg, had about four times the discharging speed of the pure salts. With prodi- gious labor the active element was extracted, in the form of the bromide, from pitchblende, and was named radium. The ore contains only 1 gram of radium per 3 tons. Properties of Radium and its Salts. The metal (at. wt. 226) is white, and turns black in the air. It is bivalent (RaBr 2 , RaCl 2 , etc.), and belongs to the calcium family. The metal and its salts show one million times the radioactivity (discharging speed) of uranium. The radiations from the salts turn oxygen into ozone, charr paper, and slowly disintegrate vessels of glass and quartz. The salts give off heat spontaneously at such a rate that 1 gram of the salt would raise a little over 1 gram of water per hour from to 100. The total heat available is calculated to be two thousand million calories per gram, or 250,000 tunes that obtainable by burning 1 gram of carbon. The Nature of the Radiations. The radiations consist of electrically charged bodies, which " ionize" the air as they shoot through it, and so render it a conductor. When a particle of a radium compound is placed in air saturated with moisture, the ionized air, cooled by expansion (p. 196), condenses the moisture to fog. Thus the track of each emitted body appears as a line of fog, and the result is a tangled mass of fog-tracks which can be observed and photographed (Wilson). There are two kinds of bodies emitted, namely positively charged atoms of the gas helium (p. 198), and particles of negative electri- city called electrons (p. 145). The atoms of helium (at. wt. 4) shoot out in almost straight 502 TEXT BOOK OF CHEMISTRY lines with a speed of about 30,000 kilometers per second, going about 3.3 cm. in air, and ploughing their way right through air- molecules by the tens of thousands before being stopped, or even diverted from a straight path. The electrons have a mass about T7W of that of a molecule of hydrogen. Being so light, their paths, straight at first, soon be- come tortuous, owing to collisions with the relatively very heavy air-molecules. The speed at first often approximates that of light (300,000 kilometers per sec.). The helium atoms can be detected also by means of Sir Wil- liam Crookes' spintharoscope (Fig. 115). The surface C is covered with zinc blende, and some of the helium atoms from the radium bromide F (at B) strike this surface and produce faint flashes of light. These are magnified by the lens A, and can be seen in a dark room after the eye has become thoroughly rested (15 to 20 minutes). The Chemistry of Radioactivity. Patient chemical study of the phenomenon has shown the nature of the changes taking place. The radium atoms are disintegrating. Each gives an atom of helium (at. wt. 4), and an atom of a monatomic gaseous element, niton (Nt, p. 254). The weight of the atom of this gas (Ramsey) is 222 (which agrees with 226 4). The gas has been liquefied (b.-p. 153). Niton is chemically inert that is, it does not combine with other elements. It disintegrates rapidly, how- ever, and in 4 days half of it has changed. Each atom loses another atom of helium and gives a radioactive solid, radium-A. This, in turn, is half gone in 4.3 minutes, giving an atom of helium and a solid, radium-B. The latter gives an electron (e) from each atom and turns into radium-C. After several further steps, which, with those described, involve altogether the loss of 5 atoms of helium and at least 5 electrons (5e) per atom, a non-radioactive substance remains. Since the combined weight of the helium is RADIUM Ra 503 5X4 = 20, and 226 - 20 = 206, this final product is found to be lead (at. wt. 207.1). Ra -> He + Nt -> He + RaA -> He + RaB -> e + RaCi -> + RaC 2 -> He + RaD 4 c + RaE -> e + RaF -> He + Pb The complex radioactivity of radium is due, therefore, to the presence with it of various amounts of all these products of its disintegration. Radioactive Change. These radioactive changes differ markedly from all the chemical changes we have hitherto dis- cussed. We can decompose a compound, or not, as we like. These changes take place of their own accord we can neither start them nor stop them. Heating greatly hastens all ordinary chemical changes. These changes proceed at the same rate at 2000 as at 200. Other chemical changes take place between atoms, these within the atoms. Why do Radium and Uranium Always Occur Together? Investigation shows that uranium, after the radium has been re- moved, is still feebly, radioactive. After a few months it re- covers its full activity. Radium can be obtained from it an indefinite number of times. Evidently the uranium atom (at. wt. 238.5) itself disintegrates, losing, by stages which have been studied, three atoms of helium and giving radium (226). Significance of Radioactivity. Uranium and thorium (also radioactive) are the elements with the highest known atomic weights. We can only imagine that these are survivors, other elements with still higher atomic weights having become extinct by disintegration. Perhaps the inactive elements now existing are simply the products of the disintegration of many radioactive elements of which the earth was once wholly composed. One interesting fact emerges from the study of radioactivity. 504 TEXT BOOK OF CHEMISTRY The Brownian movement (p. 377) has revealed to us bodies of molecular dimensions, actually behaving exactly like the mole- cules of which we had imagined matter to be composed. The study of this movement by Perrin, a French chemist, has pro- vided us with several ways of estimating the actual weights of molecules. Now, the sensitive electroscope actually enables to count single, charged atoms of helium as they enter the instrument and the fog-tracks enable us to follow their movements. The molecules and atoms of our hypotheses have thus been shown to be literal facts. Another interesting by-product of this subject is the calcula- tion that the heat given off by the disintegration of the radium known to exist in the earth (niton is found in the soil and in well waters) is sufficient alone amply to account for the maintenance of its temperature. A globe the size of the earth, possessing originally only heat energy, and cooling from a white-hot con- dition to the temperature of interstellar space, would have passed through the stage of habitable temperatures in a much shorter time than that which geological deposits and fossils show to have been actually available. The discovery of the enormous, but gradually released, disintegration energy of radium, enables us now to explain the prolonged period during which life has existed in the earth. Exercises. 1. Give a list of the classes of manganese com- pounds and, in each class, the formulae of several compounds, the prevailing color, and the valence of manganese in that class. 2. Make equations for the formation of: (a) manganous sul- phate; (b) manganous carbonate (insol.); (c) sodium manganate; (d) the oxidation of hydrochloric acid by potassium permanganate. 3. What action did we find to be catalyzed by manganese dioxide ? 4. In what respects are uranium and radium like, and in what respects unlike ordinary chemical compounds ? CHAPTER XLIII THE RECOGNITION OF SUBSTANCES, II A REVIEW OF THE METALLIC ELEMENTS As in Chapter XXXI, so here, we assume that the specimen to be identified contains a single substance. We consider now the metallic elements, and limit ourselves to the twenty-one that have been described in the context. Our review will cover, mainly, the metallurgy and some of the physical properties of each metal. Although arsenic has been stated to be a non-metallic element, and antimony to be partially non-met allic, it is more convenient in the problem of recognition to classify them with the metallic elements. They are therefore included in the twenty-one. External Examination. The color is often significant. Most of the common compounds of iron, nickel, cobalt, copper, gold, manganese and chromium are colored (see text) . A metallic luster (scrape off the tarnish) usually, though not always, indi- cates a free metal or an alloy. The crystalline form should be noted. The odor usually gives information about non-metallic constituents (p. 365) only. As regards state, the vast majority of the metals and their compounds are solids. A few are liquids (e.g. Hg and Ni(CO) 4 , p. 470) or gases. When a liquid presents itself, therefore, it is usually an aqueous solution of some com- pound. Obtain the solid by evaporation. Solubility and Reaction of the Solution. Ascertain whether the substance is soluble in water (p. 108). Note whether the solution is acid, alkaline, or neutral in reaction (p. 425). No substance can be identified by the preceding observations alone, 505 506 TEXT BOOK OF CHEMISTRY but the final conclusion as to the nature of the specimen must be in harmony with them. The Match Test (Solids).* This experiment results, in many instances, in the liberation of the metal by a metallurgical reaction. We are thus able to recognize it, largely by its phys- ical properties. A match-stick is rubbed with a warmed (and partly melted) crystal of washing soda, and is then charred in the Bunsen flame. A few particles each of the specimen and of dry sodium carbonate are made into a paste with part of a drop of water. The small mass is then stuck on the end of the match, and held in the reduc- ing region of a low Bunsen flame. The compound, whatever it may be, interacts with the sodium carbonate giving first the car- bonate and then the oxide of the metal. The oxide (if reducible) is reduced by the carbon, and the metal liberated. The end of the match is then ground in a mortar under water, and the char- coal and soda washed away. Shining or heavy grey particles indicate a metal. Results of the Match Test. The following metals, although liberated, are volatile, and are not found in the match test, namely: arsenic, antimony, bismuth, zinc, and mercury (see, therefore, p. 507). Some metals, since their oxides are not reducible by carbon in the Bunsen flame, namely calcium, magnesium, aluminium, man- ganese, and chromium, give no metal (see, therefore, p. 508). Compounds of potassium and sodium are reduced with diffi- culty by carbon, and in any case the metals are volatile, so that these metals also give no result (see p. 508). Of the twenty-one metals nine can be found by the match test : * Many experimental details, essential for the successful performance of the tests described in this chapter, are here omitted. They will be found in the Author's Laboratory Outline of Intermediate Chemistry. THE RECOGNITION OF SUBSTANCES, II, ETC. 507 Six of these metals are malleable and give shining particles. Yellow particles are gold or copper. The particles are transferred on the point of a knife-blade to a piece of filter paper, treated with a drop of nitric acid, and warmed gently. Copper dissolves leaving a green stain. Gold will not dissolve until a drop of hydrochloric acid also (making aqua regia) is added. Silver white particles may be silver, lead, tin, or platinum. Place the particles in a watch crystal, add nitric acid and warm. Silver dissolves, and the solution gives a white precipitate (AgCl) with hydrochloric acid (p. 490). Lead also dissolves, gives how- ever no precipitate (solution too dilute) with hydrochloric acid, but does give one with sulphuric acid (p. 475). Tin interacts with nitric acid, to give a white, insoluble substance (p. 478). Plat- inum is insoluble in nitric acid, but dissolves upon addition of hydrochloric acid (aqua regia). Three metals give heavy grey particles, namely, iron, cobalt, and nickel. A magnetized knife-blade attracts iron particles powerfully the others rather feebly. The particles, transferred to a piece of filter paper and touched with concentrated hydro- chloric acid, dissolve and, after warming, give a yellow stain (iron) or a blue stain (cobalt). A drop of nitric acid and warming give a greenish stain (nickel). Confirm these three by borax bead tests (pp. 468, 469, and 470). Recognition of Compounds of the Volatile Metals: Film Tests. In case of a negative result with the match test, the specimen may be a compound of one of the volatile metals (p. 506), arsenic, antimony, bismuth, mercury, or zinc. A particle of the compound is placed on the moistened end of a stiff fiber of asbestos and is held in the tip of the reducing part of a low Bunsen flame. The latter contains hydrocarbons and free hydrogen, which liber- ate the metal. A porcelain dish, held close above the end of the fiber, will condense the vapor of the metal. A scattered dust is mercury. 508 TEXT BOOK OF CHEMISTRY A continuous black metallic film is arsenic, antimony, bismuth, or zinc. Lead, if missed in the match test, being somewhat volatile, will be found here also. Dilute nitric acid dropped on the film dissolves zinc and lead instantly, the other metals slowly (Bi) or practically not at all (As,Sb). When the experiment is repeated, with the porcelain dish held high up in the flame, the vapor of the metal (except mercury) is oxidized and the oxide is deposited on the dish. A yellow de- posit is bismuth trioxide (or, if faint, PbO) . The oxides of arsenic, antimony, and zinc are white and invisible. They may be dis- tinguished, however, by blowing hydrogen iodide over the dish to produce the iodides: yellow is arsenic (Asia), orange is antimony (SbI 3 ), colorless (invisible) is zinc (ZnI 2 ). Bismuth (already identified) gives by this test a bluish-brown to red color (BiI 3 ) and lead (already identified, p. 507) a yellow (PbI 2 ). The last can be recognized because it is insoluble in water and therefore is not affected by breathing on the dish, while the other iodides are soluble and disappear (dissolve, to reappear on drying) . Recognition of Metals whose Oxides are not Reduced. There remain seven metals whose compounds give a negative result in the match and film tests. In case of such a negative result, the substance may be heated on a platinum wire in the Bunsen flame for the flame test. A brick-red color indicates a compound of calcium; a reddish violet, potassium; a very strong yellow, sodium. The borax bead test gives an amethyst color with compounds of manganese and a green with those of chromium. A small scrap of filter paper, in which a few particles of the sub- stance have been wrapped, is held by a platinum wire twisted (two or three turns) round it. The ball is charred in the Bunsen flame and is then moistened with a drop of a dilute solution of a salt of cobalt, and is heated once more. A blue color (p. 425) indi- cates a compound of aluminium (a green, a compound of zinc, p. 419) . THE RECOGNITION OF SUBSTANCES, II, ETC. 509 Compounds of the only remaining metal, magnesium, give an indefinite color (more or less pink) by the foregoing test. In case all the preceding tests have given negative results, add to a solution of the substance (1 c.c.) a few drops each of solutions of ammonium chloride, ammonium hydroxide, and sodium- hydrogen phosphate Na 2 HPO 4 . A white, crystalline precipitate (NH 4 MgP04) indicates a compound of magnesium. Recognition by Reactions in Solution. Starting with the substance in solution, its identity can also be ascertained by using reactions involving mainly precipitations and oxidations or reduc- tions. This method is of especial value when the material contains more than one substance. For example, when hydrochloric acid is added to the solution, a white precipitate (AgCl, HgCl, or PbC^) indicates the presence of silver-ion, mercurous-ion, or lead-ion. The 'other metal-ions are not precipitated, and may be separated from those three by filtration. These three chlorides can be distinguished from one another very easily. When ammonium hydroxide is added to the precipitate, silver chloride dissolves (p. 490), mercurous chloride turns black (p. 487), and lead chloride remains apparently un- changed (Pb(OH),). A definite plan of procedure enables us systematically to sepa- rate and identify every one of the positive ions. To give, and explain, the scheme in its entirety, however, would compel us to apply, and to apply rigorously and quantitatively, principles which we have touched upon, but not fully developed. Analytical chemistry is too complex a subject for brief treatment, but fasci- nating in the extreme when once it has been mastered. Exercises. 1. Name some substances that have a metallic luster, but are not metals or alloys. 2. Name the metals whose salts with active acids will give: (a) neutral aqueous solutions; (b) acid aqueous solutions. What classes of salts will give alkaline solutions? 510 TEXT BOOK OF CHEMISTRY 3. Show by what three successive changes (give equations) cupric sulphate would be acted upon in the match test (p. 506), so as finally to leave copper. 4. Give two distinctive reactions by which the green stain (p. 507) could be shown to be a salt of copper. 5. What is the action of nitric acid on tin? 6. What results do the borax bead tests give with compounds of iron, cobalt, and nickel? 7. Make equations for the action of hydrogen (p. 507) on: (a) arsenic trioxide: (b) antimony trichloride; (c) bismuth tri- chloride. 8. Make equations for the oxidation of: (a) arsenic; (b) anti- mony; (c) bismuth. 9. Make equations for the action of hydrogen iodide on the oxides of: (a) arsenic; (b) antimony; (c) zinc. 10. What are the interactions of a salt of cobalt heated (dry) with a compound of: (a) aluminium; (b) zinc? 11. Make an equation for the interaction of ammonium hy- droxide, magnesium sulphate and sodium phosphate. 12. Make equations for the actions of ammonium hydroxide solution on chlorides of: (a) silver-ion; (b) mercurous-ion; (c) lead-ion. APPENDIX I. The Metric System Length. 1 meter = 10 decimeters = 100 centimeters (100 cm.) = 1000 millimeters (1000 mm.). 1 decimeter = 10 centimeters = nearly 4 inches. Volume. 1 liter = 1000 cubic centimeters (1000 c.c.) = a cube 10 cm. X 10 cm. X 10 cm. 1 liter = 1.057 quarts (U.S.) or 1.136 quarts (Brit.). 1 fl. ounce (U.S.) = 29.57 c.c. 1 ounce (Brit.) = 28.4 c.c. Weight. 1 gram (g.) = wt. of 1 c.c. of water at 4 C. 1 kilo- gram = 1000 g. 1 kilogram = 2.2 Ibs. avoird. (U.S. and Brit.). 1 ounce avoird. (U.S. and Brit.) = 28.35 g. 1 nickel (U.S.) weighs 5 g. 1 halfpenny (Brit.) weighs 5 to 5.7 g. II. Scale of Hardness Each of the following minerals will scratch the surface of a specimen of any one preceding it in the list. 1. Talc 6. Felspar 2. Gypsum (or NaCl) 7. Quartz 3. Calcite (or Cu) 8. Topaz 4. Fluorite 9. Corundum 5. Apatite 10. Diamond Glass is slightly scratched by 5, and easily by those following. Glass will not scratch 5 distinctly, but will scratch those pre- ceding 5. A good knife scratches 6 slightly, but not those following. A file will scratch 7, but not those following. 511 512 TEXT BOOK OF CHEMISTRY III. Temperatures Centigrade and Fahrenheit Upon the centigrade scale, the freezing-point of water is C. and the boiling-point 100 C. Upon the Fahrenheit scale, the same points are 32 F. and 212 F., respectively. The same in- terval is thus 100 on the one scale and 180 on the other. The degree Fahrenheit is therefore or - of 1 Centigrade. Any y temperatures can be converted by using the formulae: C. = ^ (F. - 32), F. = - (C.) + 32. y 5 The following table (IV) contains the temperatures from C. to 35 C., with the corresponding values on the Fahrenheit scale (32 F. to 95 F.). IV. Vapor Pressures of Water Both the Fahrenheit (F) or ordinary and the Centigrade (C) temperatures are given. Temperature. Pressure, mm. Temperature. Pressure, mm. F. C. F. C. 32 4.6 71.6 22 19.7 41 5 6.5 73.4 23 20.9 46.4 8 8.0 75.2 24 22.2 48.2 9 8.6 77.0 25 23.6 50.0 10 9.2 78.8 26 25.1 51.8 11 9.8 80.6 27 26.5 53.6 12 10.5 82.4 28 28.1 55.4 13 11.2 84.2 29 29.8 57.2 14 11.9 86.0 30 31.5 59.0 15 12.7 87.8 31 33.4 60.8 16 13.5 89.6 32 35.4 62.6 17 14.4 91.4 33 37.4 64.4 18 15.4 93.2 34 39.6 66.2 19 16.3 95.0 35 41.8 68.0 20 17.4 69.8 21 18.5 212.0 100 760.0 APPENDIX 513 V. Order of Activity of the Metals Each metal displaces from combination those following it in the list (see p. 54). Potassium Sodium Calcium Magnesium Aluminium Manganese Zinc Chromium Iron Nickel Tin Lead Hydrogen Copper Bismuth Antimony Mercury Silver Platinum Gold INDEX Acids are all listed under " acid," and salts under the positive radical. ACETYLENE, 209, 211 torch, 211 Acid, acetic, 226 boric, 348 carbonic, 201 hydrochloric, 119 chemical properties, 124 composition, 140 molecular view, 121 physical properties, 123 preparation, 119 uses, 127 hypochlorous, 304 nitric, 313 nitrous, 320 orthophosphoric, 338 sulphuric, chamber process, 280 chemical properties, 285 contact process. 278 physical properties, 285 uses, 287 sulphurous, 276 Acids, 163 bases, salts, common properties, 166 of sulphur, 273 organic, 371 properties, 178 Agriculture, 447-454 Air, 26 Alcohol, denatured, 225 Alkalies and bases, 162 Aluminates, 425 Aluminium, 422 hydroxide, 424 Aluminium, oxide, 425 sulphate, 426 uses, 423 Alums, 426 Ammonia, 256 chemical properties, 259 physical properties, 258 uses of, 261 Ammonium hydroxide, 260 salts, 260 Antimony, 341 alloys, 342 Appendix, 511 Aqueous vapor, tension of, 46 Argon, 253 Arsenic, 340 trioxide, 341 Atmosphere, 247 dust, 251 humidity, 248 Atomic numbers, Moseley's, 296 weights, 79, 86 Attributes, 7 BAKELITE, 386 Baking powders, 353 Bases, 163 properties, 179 Battery, storage, 476 Beer, 225 Bismuth, 342 Bleaching, 307 powder, 307 Blue-prints, 467 Bodies, 6 515 516 INDEX Boiling-point, acids, bases, salts, 172 Borax, 349 Boron, 348 Bromine, 324 CALCIUM, 387 carbonate, 388 chloride, 404 cyanamide, 405 hydroxide, 390 oxalate, 391 oxide, 389 phosphates, 338 sulphate, 399 Calculations, 128, 142 Carbides, 192 Carbon, 188 amorphous, 190 chemical properties, 191 dioxide, 198 chemical properties, 200 physical properties, 200 plant food, 204 disulphide, 271 monoxide, 202 chemical properties, 203 physical properties, 203 poison, 204 tetrachloride, 192 Catalytic action, 32 Celluloid, 384 Cellulose, 227 Cement, 428 Charcoal, animal, 445 wood, 436 Chemical change, and energy, 147 varieties, 126 Chemical equilibrium, 228 properties, 33 units of weight, 73 Chlorine, chemical properties, 138 physical properties, 137 preparation, 133 Chromium, 498 Chromic hydroxide, 500 oxide, 500 sulphate, 499 Clay and pottery, 427 Coal, 436 analysis 437 as fuel, 439 gas, 437 Cobalt, 469 salts of, 469 Coke ovens, 438 uses, 439 Collodion, 385 adsorption, 380 suspension, 377 Colloids, properties, 378 Combining proportions, measure- ment, 36 Combustion, spontaneous, 41 Components, 6 Conditions, 7 Constituents, 8 Copper, 480 plating, 483 refining, 484 salts, 483 uses, 481 Crystallization, 110 Crystal structure, 298 Cupric oxide, 483 sulphate, 482 Cuprous oxide, 482 DECOMPOSITION, 16 Diamonds, 188, 189 Diet, normal, 444 Diffusion, 92 Dissociation, 104 Dyeing, 429 ELEMENTS, 17 list of common ones, 19 Electric furnace, 193 Electrolysis, 55 INDEX 517 Electrolytes, conductivity of, 167 Electrotypes, 484 Energy, 149 and chemical action, 311 and chemical activity, 153 in chemistry, 152 interconvertibility, 151 Enzymes, 223 Equations, 81 making of, 100 balancing, 101 ionic, 175 Equilibrium, 64 Equivalent weights, 53 Esters, 372 Evaporation, 14 Explanation, 22 FATS and oils, 373 Fats, chemical properties, 374 digestion, 443 Fermentation, alcoholic, 223 Ferric chloride, 464 hydroxide, 464 Ferrous carbonate, 466 chloride, 46 hydroxide, 465 sulphate, 464 Fertilizers, 433 nitrogen, 434 Filtration, 13 Fixation of nitrogen, 319 Flame, 214-217 Flour, 5 Fluorine, 331 Food, 440-442 fuel value, 445 Formulae, 80 making of, 98 Freezing-point, acids, bases, salts, 172 Fuels, 435 GASES, diatomic, 77 liquefaction of, 92 Gases, measurement of, 44 summary, 94 water, 202 Gasoline, 213 Glass, 346 Glucose, from starch, 220 Gluten, 5 Gold, 492 properties, 493 uses, 493 Granite, 4 Graphite, 189, 190 Guncotton, 383 Gunpowder, 360 HALOGENS as a family, 333 Hard water, 400 Helium, 254 Hydrocarbons, 208 chemical properties, 212 Hydrolysis of salts, 356 Hydrogen bromide, 325 Hydrogen, chemical properties, 56 fluoride, 332 iodide, 329 peroxide, 302 physical properties, 55 preparation, 50 sulphide, 267 IMPURITIES, 6 Ink, 467 Iodine, 327 Ion-product constancy, 412 Ionic equilibrium, 406 lonization, and chemical activity, 171 reversible, 174 Ions, and conductivity, 173 and displacement, 178 and double decomposition, 177 and radicals, 173 carry electrical charges, 174 Iron, 455 Bessemer process, 460 518 INDEX [ron, cast, 457 open-hearth, 460 other salts, 468 pure, 463 wrought, 457 KINDLING conditions, 39 Krypton, 254 LAW, Avogadro's, 83, 92 Boyle's, 90 Charles', 90 chemical change, 7 combining weights, 78 component substances, 5 definite proportions, 19 definition, 8, 22 Dulong and Petit 's, 87 Gay-Lussac's, 61 Le Chatelier's, 242 of effect of heating, 27 Van't Hoff's, 241 Lead, 472 acetate, 474 oxides of, 473 salts, 475 use of, 473 white, 474 Liquids, molecular relations, 95 MAGNESIUM, 415 hydroxide, 416 oxide, 416 salts of, 416 Manganese, 496 dioxide, 497 Manganous salts, 497 Mass, conservation of, 20 Matches, 337 Matter, states of, 65 structure and behavior, 88 Mercuric oxide, 15 Mercury, 485 salts, 486 Mercury, uses, 486 Metals, order of activity, 54 Methane, 210 Mixtures, 6, 12 Molecular hypothesis, 47 Molecular theory, history, 95 Molecular weights, 84 Mortar, 391 NEON, 254 Nickel, 469 plating, 485 salts of, 470 uses, 470 Niton, 254 Nitric oxide, 317, 318 Nitrogen, 245 chemical properties, 246 peroxide, 315, 318 physical properties, 246 Nitroglycerine, 382 Nitrous oxide, 321 OIL refining, 209 Oleomargarine, 374 Organic analysis, 483 Oxidation, 40 Oxygen, 27 chemical properties, 34 history of, 28 physical properties, 33 preparation, 29, 30, 31 uses of, 41 Ozone, 300 PAINTS, 475 Paper manufacture, 277 Periodic system, 288 Petroleum, 209 Phosphorus, manufacture, 336 Photography, 490 Physical properties, 33 Physics in chemistry, 21 Plant nutrition, 433 INDEX 519 Plants, 432 Plaster of Paris, 399 Platinum, 493 uses, 494 Potassium, 357 bicarbonate, 359 carbonate, 358 chloride, 358 chromate, 498 dichromate, 499 ferricyanide, 467 ferrocyanide, 466 hydroxide, 359 manganate, 497 nitrate, 359 permanganate, 497 Precipitation, 125 theory of, 391 Properties, specific, 7 Proteins, digestion, 443 RADICALS, If Radioactivity, 502 meaning of, 503 Radium, 500 Recognition of substances, metallic, 505 non-metallic, 362 Reduction, 57 Relations of liquid and vapor, 63 Reversible chemical actions, 70 Rouge, 465 Rusting, 9 SALTS, 164, 179 Saturated solutions, 410 Saturation, 108 Scientific method, 22 Silicon, 345 Silk, artificial, 385 Silver, 488 halides, 490 nitrate, 490 properties, 489 Silver, uses, 489 Smokeless powder, 384 Soap, chemical properties, 376 cleansing power, 379 colloidal, 378 emulsifying action, 379 manufacture, 37J removes dirt, 380 Sodium, 158 bicarbonate, 352 carbonate, 355 chloride ' 161 '61 ^ns, 96 wature, 109 heinical, 117 saturated, 116 Solutions, boiling-points, 114 freezing-points, 114 molar, 116 normal, 115 vapor pressure of, 113 Solvents, 108 influence of solute upon, 112 Speed of chemical actions, 154 Spintharoscope, 502 Stains, removing, 468 Starch, 219 digestion, 442 Steam, 62 Steel alloys, 462 crucible, 458 tempering, 459 Substance, 3 Substances, compound, 16 simple, 16 Sugars, 221 Sugar, beet sugar, 222 \ne-sugar, 221 222 dies, 287 Sulphide ;, 271 u chemical properties, 266 *o^ 274 physical properties, 264 78 a, 111 v M i. Tin, * s salts, uses, 4/o ULTRAMARINE, VALENCE, 181- Ventilation, 250 Volume, standard (22.4 liters), 74 , 105 Water as solvent, 06 Water, chemical properties, 67 coagulation process, 427 composition, 59 composition by volume, 60 physical properties, 61 purification, 66 Water gaa, carburetfced, 213 Waters, natural, 66 Whiskv, 225 UNIVERSITY OF CALIFORNIA LIBRARY 2 a I o slJ i JS |1| g< fl. -8 W ** W H w 1 c? w w B H s a O OQ w P w w ffl S5 S a" 6 8 3. i YC 22085 INTERNATIONAL ATOMIC WEIGHTS (1917) Aluminium Al 27 , 1 Antimony Sb 120.2 Argon A 39.88 Arsenic As 74.96 Barium Ba 137.37 Bismuth Bi 208.0 Boron B 11.0 Bromine Br 79.92 Cadmium Cd 112.40 Caesium Cs 132.81 Calcium Ca 40.07 Carbon C 12.005 Cerium Ce 140.25 Chlorine Cl 35.46 Chromium Cr 52.0 Cobalt Co 58.97 Columbium Cb 93 . 1 Copper Cu 63.57 Dysprosium Dy 162.5 Erbium Er 167.7 Europium Eu 152.0 Fluorine F 19.0 Gadolinium Gd 157.3 Gallium Ga 69 . 9 Germanium Ge 72 . 5 Glucinum Gl 9.1 Gold Au 197.2 Helium He 4.00 Holmium Ho 163.5 Hydrogen H 1 .008 Indium In 114.8 Iodine I 126.92 Iridium Ir 193.1 Iron Fe 55.84 Krypton Kr 82.92 Lanthanum La 139.0 Lead Pb .207.20 Lithium Li 6.94 Lutecium Lu 175 .0 Magnesium Mg 24.32 Manganese Mn 54.93 Mercury Hg 200.6 = 16. Molybdenum Mo 96.0 Neodymium Nd 144.3 Neon Ne 20.2 Nickel .Ni 58.68 Niton (radium emanation) Nt 222.4 Nitrogen N 14.01 Osmium Os 190.9 Oxygen O 16.00 Palladium Pd 106.7 Phosphorus P 31.04 Platinum Pt 195.2 Potassium K 39.10 Praseodymium Pr 140.9 Radium Ra 226.0 Rhodium Rh 102.9 Rubidium Rb 85.45 Ruthenium Ru 101 .7 Samarium Sa 150 . 4 Scandium Sc 44. 1 Selenium Se 79.2 Silicon Si 28.3 Silver Ag 107.88 Sodium Na 23.00 Strontium Sr 87.63 Sulphur S 32.06 Tantalum Ta 181 . 5 Tellurium Te 127.5 Terbium Tb 159.2 Thallium Tl 204.0 Thorium Th 232.4 Thulium Tm 168.5 Tin Sn 118.7 Titanium Ti 48. 1 Tungsten W 184.0 Uranium U 238.2 Vanadium V 51.0 Xenon Xe 130.2 Ytterbium (Neoytterbium) Yb 173.5 Yttrium Yt 88.7 Zinc Zn 65.37 Zirconium Zr 90 . 6