Jleto $ork g>tate College of Agriculture &t Cornell ®mbers CONTENTS. droxide — Cuprous Oxide — Cupric Hydroxide — Cupric Oxide — Cupric Sulphate — Applications — Cupric Nitrate — Cupric Arsenite — Cupric Carbonates — Cuprous Sulphide — Cupric Sulphide — Copper-plating — Reactions which are of Special Value in Chemical Analysis — Silver — General — Forms in which Silver occurs in Nature — Metallurgy of Silver — Puri- fication — Properties — Allotropic Forms of Silver — Alloys of Silver — Argentous Chloride — Silver Chloride, Argentic Chloride — Silver Bromide and Iodide — Application of the Chloride, Bromide, and Iodide of Silver in the Art of Pho- tography — Silver Triazoate — Silver Oxide — Other Oxides of Silver — Sulphides of Silver — Silver Nitrate, Argentic Ni- trate — Silver Cyanide — Reactions which are of Special Value in Chemical Analysis — Gold — General — Forms in which Gold occurs in Nature — Metallurgy of Gold — Properties — Alloys of Gold— Chlorides of Gold— Gold Sulphide. Experiments : Copper and its Salts— Silver and its Salts, . . 537 CHAPTEE XXIX. ELEMENTS OF FAMILY II, GROUP BI ZINC — CADMIUM — MERCURY. General — Zinc — General — Forms in which it occurs in Nature — Metallurgy — Properties — Applications — Alloys — Zinc Chlo- ride — Zinc Hydroxide — Zinc Oxide — Zinc Sulphide — Zinc Sulphate — Zinc Carbonate — Reactions which are of Special Value in Chemical Analysis — Cadmium — General — Prepa- ration and Properties — Cadmium Sulphide — Analytical Reac- tions — Mercury — General — Forms in which Mercury occurs in Nature — Metallurgy of Mercury — Properties — Applica- tions — Amalgams — Mercurous Chloride — Mercuric Chloride, or Corrosive Sublimate — Mercurous Iodide — Mercuric Iodide — Mercuric Oxide — Mercuric Sulphide — Mercuric Cyanide — Mercurous Nitrate— Mercuric Nitrate — Compounds formed by Salts of Mercury with Ammonia — Reactions which are of Special Value in Chemical Analysis. ELEMENTS OF FAMILY III, GROUP B ! GALLIUM— INDIUM — THALLIUK. General — Gallium — Indium — Thallium. Experiments : Zinc and its Salts — Mercury and its Salts, . . 559 CONTENTS. xvn CHAPTER XXX. ELEMENTS OP FAMILY IV, GROUP B '. GERMANIUM — TIN — LEAD. PAGE General — Germanium — Tin — General — Occurrence — Metallurgy — Properties — Applications — Alloys — Stannous Chloride — Stannic Chloride — Stannic Hydroxide — Stannic Oxide — Stannous Sulphide — Stannic Sulphide — Reactions which are of Special Value in Chemical Analysis — Lead — General — Forms in which Lead occurs in Nature — Metallurgy — Prop- erties — Applications — Lead Chloride — Lead Tetrachloride — Lead Iodide — Oxides of Lead — Lead Suboxide — Lead Oxide — Lead Peroxide — Storage Battery — Lead Sulphide — Lead Nitrate — Lead Carbonate — Lead Sulphate — Reactions which are of Special Value in Chemtcal Analysis — Lanthanum — Cerium — Didymium — Praseodymium and Neodymium. Experiments : Tin and its Compounds — Lead and its Com- pounds, . 576 CHAPTER XXXI. ELEMENTS OP PAMILT VI, GROUP A : CHROMIUM — MOLYBDENUM — TUNGSTEN — URANIUM. General — Chromium — General — Forms in which Chromium oc- curs in Nature — Preparation — Properties — Chromic Chloride — Chromic Hydroxide — Chromic Oxide — Chromic Sulphate — Chrome-Alums — Chromic Acid and the Chromates — Potas- sium Chromate — Potassium Bichromate — Chromium Trioxide — Relations between theChromates and Bichromates — Sodium Chromate and Sodium Bichromate — Barium Chromate — Lead Chromate — Reactions which are of Special Value in Chem- ical Analysis — Molybdenum — General — Occurrence and Prep- aration — Molybdic Acid and the Molybdates — Lead Molyb- date — Phospho-molybdic Acid — Tungsten — General — Occur- rence and Preparation — Properties — Tungstic Acid and the Tungstates — Uranium — General — Occurrence and Prepara- tion—Properties — Oxides — Uranous Salts— TJranyl Salts — Uranates. Experiments : Chromic Acid and the Chromates, 593 CONTENTS. CHAPTER XXXII. ELEMENTS OF FAMILY Til, GROUP A : MANGANESE. P General — Forms in which Manganese occurs in Nature — Prepar- ation and Properties — Manganous Chloride — General Re- marks concerning the Oxides — Manganous Hydroxide — Manganous-manganic Oxide — Manganic Oxide — Manganese Dioxide — Weldon's Process for the Regeneration of Man- ganese Dioxide in the Preparation of Chlorine — Sulphides — Manganous Sulphate — Manganic Sulphate — Manganic Acid and the Manganates — Permanganic Acid and the Perman- ganates — Potassium Permanganate — Reactions which are of Special Value, in Chemical Analysis. ExfttujiLHiNTS : Manganese and its Compounds, CHAPTER XXXIII. ELEMENTS OF FAMILY VIII, SUB-GROUP A. IKON — COBALT — NICKEL. General — Iron — Introductory — Forms in which Iron occurs in Nature — Metallurgy — Varieties of Iron — Properties of Iron — Ferrous Chloride — Ferric Chloride — Cyanides — Potassium Ferrocyanide — Ferrohy drocyanic Acid — Ferric Ferrocyanide, ' or Prussian Blue — Potassium Ferricyanide — Ferrihydro- cyanic Acid — Ferrous Ferricyanide — Ferrous Hydroxide — Ferric Hydroxide — Ferrous-Ferric Oxide — Soluble Ferric Hydroxide — Ferric Oxide — Ferrous Sulphide — Ferrous Car- bonate — Ferrous Sulphate — Ferric Sulphate — Ferrous Phos- phate — Iron Disulphide — Reactions which are of Special Value in Chemical Analysis — Cobalt — General — Occurrence and Preparation — Properties — Cobaltous Chloride — Cobalt- ous Hydroxide — Cobaltic Hydoxide — Cobalt Sulphide — Cyanides — Smalt — Nickel — General — Occurrence and Prep- aration — Properties — Alloys — Other Applications of Nickel — Nickelous Hydroxide — Nickelic Hydroxide — Cyanides — Nickel Carbonyl — Reactions of Cobalt and Nickel which are of Special Value in Chemical Analysis. Experiments : Iron and its Compounds 620 CONTENTS. CHAPTER- XXXIV. ELEMENTS OF FAMILY VIII, SUB-GROUP B : RUTHENIUM — RHODIUM — PALLADIUM. ELEMENTS OF FAMILY VIII, SUB-GROUP C : OSMIUM — IRIDIUM — PLATINUM. PAGE General — The Platinum Metals — Metallurgy — Ruthenium — Prep- aration — Properties — Osmium — Preparation — Properties — Rhodium — Iridium — Preparation — Properties — Palladium — Preparation — Properties — Palladium-Hydrogen — Platinum — Preparation — Properties — Applications of Platinum — Alloys of Platinum — Chlorides — Chlorplatinio Acid — Sulphides. Experiments : Platinum — Chemical Analysis 641 CHAPTER XXXV. SOME FAMILIAR COMPOUNDS OF CARBON. Organic Chemistry — Occurrence of the Compounds of Carbon — Formation of Hydrocarbons — Distillation of Coal — Distilla- tion of Wood — Distillation of Bones — Fermentation — Classes of Compounds of Carbon — Compounds of Carbon and Hydrogen — Petroleum — Refining of Petroleum — Hydrocar- bons in Petroleum — Homology — The Ethylene Series of Hydrocarbons — The Acetylene Series— The Benzene Series — Marsh-gas, Methane, Fire-damp — Substitution-products of the Hydrocarbons — Chloroform — Iodoform — Ethylene, Ole- fiant Gas — Acetylene — Methyl Alcohol, Wood-spirit — Ethyl Alcohol, Spirits of Wine — What Change takes Place in Sugar ? — What Causes the Change ? — Germs in the Air — Different Kinds of Fermentation — Distillation of Fermented Liquids — Properties of Alcohol — Uses of Alcohol — Glycerin — Properties — Acetic Aldehyde, Ordinary Aldehyde — Chloral — Formic Acid — Acetic Acid — Properties — Uses — Salts of Acetic Acid — Fatty Acids — Butyric Acid — Palmitic Acid— Stearic Acid — Soaps — Use of Soap — Action of Soap on Hard Waters— Relations of the Soap Industry to Other Industries — Oxalic Acid — Lactic Acid — Malic Acid — Tartaric Acid — Citric Acid — Ether — Action of Acids upon Alcohols — Saponification — Fats — Butter — Ethereal Salts as Essences — Nitroglycerin Comparison of Formulas — Alcohols — More Complex Alcohols — Radicals or Residues — Acids. Experiments : Fermentation — Aldehyde — Soap — Hard Water, . 649 xx CONTENTS. CHAPTER XXXVI. OTHER COMPOUNDS OF CARBON. PAGE The Carbohydrates — Grape-sugar, Glucose, Dextrose — Formation of Dextrose — Manufacture of Dextrose or Glucose — Proper- ties — Levulose, Fruit-sugar — Cane-sugar — Sugar-refining — Molasses — Properties of Sugar — Sugar of Milk, Lactose — Souring of Milk — Cellulose — Properties — Gun-cotton, Py- roxylin, Nitrocellulose — Collodion — Celluloid — Paper — Starch — Manufacture of Starch — Properties — Flour — Bread- making — Aromatic Compounds — Nitrobenzene — Aniline — Aniline Dyes — Phenol, Carbolic Acid — Oil of Bitter Almonds, Benzoic Aldehyde — Benzoic Acid — Balsams and Odoriferous Resins— Gallic Acid — Tannic Acid, Tannin — Tanning — In- digo — -Naphthalene — Anthracene — Alizarin — Glucosides — Myronic Acid — Alkaloids — Quinine — Cocaine — Nicotine — Morphine and Narcotine, 667 Index 677 A COLLEGE TEXT-BOOK OF CHEMISTRY. CHAPTEE I. CHEMICAL AND PHYSICAL CHANGE. — EARLIEST CHEMICAL KNOWLEDGE.— LAW OF THE INDE- STRUCTIBILITY OF MATTER.— LAW OF DEFINITE PROPORTIONS.— LAW OF MULTIPLE PROPORTIONS. —THE ELEMENTS. Matter and Energy. — The sensible universe is made up of matter and energy. It is difficult to give satisfactory definitions of either of these terms, but, in a general way, it may be said that matter is anything which occupies space, and energy is that which causes change in matter. It requires but little observation to show that there are many' kinds of matter, and apparently many kinds of energy. As examples of the different kinds of matter we have the many varieties of rocks and earth, as granite, limestone, quartz, clay, sand, etc.; the plants and their fruits; the substances which enter into the composition of animals; and innumerable manufactured products. As examples of the different forms of energy, we have heat, light, motion, etc. Under the influence of the forms of energy the forms of matter are constantly undergoing change. Everywhere these changes are taking place. Changes in position and in temperature appeal most directly to our senses, and are most easily studied. But there are many other kinds of change which are of the highest importance. Thus there are electrical changes, manifestations of which we see in thunder-storms; there 2 COLLEGE CHEMISTRY. are magnetic changes which may be studied to some extent by means of the magnetic needle ; and there are, further, what are called chemical changes which affect the compo- sition of substances. Chemical Change. — For the purpose of study it is con- venient to distinguish between two classes of changes in matter, the difference between which can best be made clear by means of examples. Consider the changes in- cluded under the head of fire. We see substances destroyed by fire, as we say. They disappear as far as we can deter- mine by ordinary observation. When iron is exposed to the air a serious change takes place. It becomes covered with a reddish-brown substance which we call rust. If the piece of iron is comparatively thin, and it is allowed to lie in the air long enough, it is completely changed to the reddish-brown substance, and no iron as such is left. If the juices from fruits, as from apples, are allowed to stand in the air, they undergo change, becoming sour, and a somewhat similar change takes place in milk. If a spark is brought in contact with gunpowder there is a flash and the powder disappears, a dense cloud appearing in its place. In the changes referred to the substances changed dis- appear as such. After the fire, the wood or the coal, or whatever may be burned, is no longer to be found. The rusted iron is no longer iron. The gunpowder after the flash is no longer gunpowder. Changes of this kind in which the substances disappear and something else is formed in their place are known as chemical changes. Physical Change.— There are many changes taking place that do not affect the composition of substances. Iron, for example, may be changed in many ways and still remain iron. It may become hotter or colder. Its posi- tion may be changed, or, as we say, it may be moved. The iron may be struck in such a way as to give forth sound. It" may be made so hot that it gives light. When, for example, it becomes red-hot, it can be seen in a dark room. A piece of iron may be changed further by con- PHYSICS AND CHEMISTRY. 3 necting it with what is known as a galvanic battery. A current of electricity then passes through it, and we can easily recognize the difference between a piece of iron through which a current of electricity is passing and one through which no current is passing. The former when brought into certain liquids will at once change their com- position, while the latter will not cause such change. Finally, when a piece of iron is brought in contact with loadstone, it acquires new properties. It now has the power to attract and hold to itself other pieces of iron. In all these cases, the iron is changed, but it remains iron. After the moving iron comes to rest, it is exactly the same thing that it was before it was moved. After the iron that is giving forth sound has ceased to give forth sound, it returns to its original condition. Let the heated iron alone, and it cools down, soon ceasing to give light, and presenting no evidence of being warm. Eemove the iron from contact with the galvanic battery, and it loses those properties which are due to the current of electricity. In time, the iron which is magnetized by contact with the loadstone loses its magnetic properties. It then no longer has the power to attract other pieces of iron, and does not differ from ordinary iron. While iron has been taken as an example, other sub- stances undergo similar changes. These changes which do not affect the composition of the substances are called physical changes. Physics and Chemistry. — According to what has been said, we have two classes of changes presented to us for study: (1) Those which do not affect the composition of sub- stances, or physical changes. (2) Those which do affect the composition of substances, or chemical changes. That branch of science which has to deal with physical changes is known as Physics. And that which has to deal with chemical changes is known as Chemistry. Everything that has to do with motion, heat, light, 4 COLLEGE CHEMISTRY. sound, electricity, and magnetism is studied under the head of Physics. Everything that has to do with the composition of substances is studied under the head of Chemistry. It is, however, impossible to study these two subjects entirely independently of each other. Whenever a chemical change takes place, it is accompanied by physical changes; and in order that the former may be clearly understood, a study of the latter is necessary. Earliest Chemical Knowledge. — Those substances which are most abundant and most widely distributed in nature were, of course, the first known and studied; and the same is true of those chemical changes which occur most com- monly and produce the most striking eftects. Simply by observing those things which surround us and those changes in composition which take place naturally, a con- siderable amount of chemical knowledge might be gained, and indeed the earliest knowledge of chemistry was acquired in this way. It was not, however, until men came to experiment upon the substances which they found in nature, that knowledge of chemical changes made rapid progress. Since then an enormous amount of knowledge has been gained, and every year the stock is increased by new discoveries, until the field appears almost boundless. Law of the Indestructibility of Matter. — One of the first facts of fundamental importance established by the study of chemical phenomena is that after a chemical change has taken place there is just as much matter as there was before the change. Whenever matter apparently disappears, it continues to exist in some other form. If it were possible to annihilate matter or to call it into being, it would be of little or no value to weigh things. The first fundamental law bearing upon the changes in com- position which the different forms of matter undergo is the law of the indestructibility of matter. While, if we think of it, we can not conceive that this great law should not be true, we must not forget that the only way in which its truth could be established was by experiment. The law may be stated thus : f CONSERVATION OF ENERGY. 5 Whenever a change in the composition of substances takes place the amount of matter after the change is the same as before the change. According to this, and assuming that the law has always held good, it follows that the amount of matter in the uni- verse is the same to-day as it has been from the beginning. Transformations are constantly taking place, but these involve no increase nor decrease in the total amount of matter. Conservation of Energy. — Just as matter is neither created nor destroyed, so it has been made probable that the total amount of energy is unchangeable. One of the greatest discoveries in science was the recognition of the fact that one form of energy can be transformed into others, and that in these transformations nothing is lost. We now know that for a certain amount of heat we can get a certain amount of motion, and that for a certain amount of motion we can get a certain amount of heat. We know that a similar definite relation exists between heat and electrical energy, and between these and chemical energy. We know, for example, that a definite amount of heat can be obtained by burning a definite amount of a given sub- stance, and we know also that with a definite amount of heat we can produce a definite amount of chemical change. Modern investigation has shown that all the different forms of energy are convertible one into the other without loss. This important fact is generally spoken of as the law of the conservation of energy. Transformations of energy are taking place constantly, as transformations of matter are, but the total amount in each case remains the same. Composition of Matter. — The fact that first impresses one in studying the various forms of matter found in the earth is their great variety. We find an almost infinite number of kinds of matter, and the question at once sug- gests itself, of what are these things composed ? This question has long been asked, and it will be long before an entirely satisfactory answer is reached. Still, much 6 COLLEGE CHEMISTRY. more is now known in regard to the subject than was known in past ages, and some progress is constantly being made towards a solution of the problem. Elements. — As chemical experimenting upon things advanced, the fact impressed itself more and more strongly upon investigators, that, of the large number of substances known, some can be converted into simpler ones by chem- ical action and some cannot. In other words, some sub- stances like water can be broken down by various methods into two or more others of different properties, and these . when brought together again under proper conditions form the original substance. Water can be decomposed into two gases, hydrogen and oxygen. Elaborate experiments have shown that the weight of water decomposed is exactly equal to the weight of the hydrogen plus that of the oxygen obtained, and that when the hydrogen and oxygen are brought together again under proper conditions exactly as much water is formed as was originally decomposed. It appears, therefore, that water consists of at least .two simpler substances. A similar conclusion is reached by a study of by far the largest number of the substances with which we have to deal. On the other hand, no treatment to which hydrogen and oxygen have been subjected has, as yet, effected their decomposition. They can be made to combine with other substances, as, for example, with each other, and thus form more complex substances, but nothing simpler than hydrogen has ever been obtained from hydrogen, and nothing simpler than oxygen has ever been obtained from oxygen. Whether the decomposition of these substances will ever be effected is a question that cannot be answered. All that we know is' that at present they cannot be decomposed. We therefore speak of them as elements, meaning by the term, that, with the means now at the disposal of chemists, it is impossible to get simpler substances from them. There are at present seventy-six substances known which are called elements for the same reasons that hydrogen and oxygen are called elements. It is quite possible that the number may be CHEMICAL ACTION. 7 increased in the future, and it is also quite possible that the number may be decreased. New elements will in all probability be discovered, and probably some of the sub- stances now included in the list of elements will eventually be shown to be capable of decomposition. The view at present held in regard to the forms of matter which go to make up that part of the universe which comes under our observation is that they are all composed of the elementary substances. Many of them, like water, are composed of only two elements; others of three; and still others of four, five, six, and more; but most of them are comparatively simple, and rarely does any one contain more than four or five elements. Of the seventy elements known, only about twelve enter into the composition of most things with which we com- monly have to deal. The others occur in relatively small quantity. Chemical Action. — In the last paragraph it was stated that- most substances can be decomposed, and that under proper conditions the elements combine. We must now inquire more carefully into the meaning of these expres- sions. Among the elements are the well-known substances lead, iron, and sulphur. If some finely-divided iron is brought in contact with sulphur, apparently no action takes place. If the two are put in a mortar and mixed, no matter how thoroughly, there is no evidence of action. The mixture has, to be sure, a different appearance from that of either constituent, but still both substances are present, and can be recognized in various ways. If, for example, a little of the mixture is examined with the aid of the microscope, particles of iron and of sulphur will be recognized lying side by side. If, further, the mixture is treated with the liquid, carbon disulphide, which has the power to dissolve the sulphur but not the iron, the sulphur will be dissolved while the iron will be left unchanged. Finally, if a dry magnet is introduced into the mixture, the iron will adhere to it, and by careful manipulation the two constituents can be separated. These facts furnish 8 COLLEGE CHEMISTRY. evidence that both iron and sulphur are present in the mixture in unchanged condition, just as sugar and sand are present in a mixture of these two substances. If now the mixture of sulphur and iron is heated in a dry test- tube, marked changes will take place, and there will be formed a black substance entirely different from either of the elements employed in the experiment. Carbon disul- phide can no longer extract sulphur from it. The magnet can no longer pick out the iron, and under the microscope one homogeneous substance is seen instead of the two ele- mentary substances. If the experiment is performed with proper precautions, the amount of matter after the action will be found to be exactly the same as before the action. A serious change has taken place, but no change in the amount of matter. The act is one of chemical combina- tion, and the substance formed is called a chemical corn- pound. A few other examples will aid in making the conception of chemical combination clear. When a bit of phosphorus is brought in contact with a little iodine, action takes place at once; the two elements combine, losing their own characteristic properties and forming a compound with properties quite different from those of the constituents. When the gases hydrogen and oxygen are brought together and a spark is passed through the mixture, an explosion occurs, and, in place of the gases, the liquid, water, is formed. When sulphur burns in the air the product formed is a pungent gas. It has been shown that the act consists in the combination of the sul- phur with the gas, oxygen, which is contained in the air. All these cases are examples of chemical combination. But chemical action may be of the opposite kind, that is to say, instead of being combination, it may be decomposi-. tion. Thus, water which is formed by the chemical com- bination of hydrogen and oxygen may, by proper methods, be decomposed into the same elements. We may con- veniently think of that which causes elements to combine as an attractive force exerted between the elements. Now, when some power which can overcome this attraction is CHEMICAL AFFINITY. $ brought to bear upon a compound, decomposition takes place, and the elements are, as we -say, set free. When, for example, the substance known as red oxide of mercury or mercuric oxide is heated to a sufficiently high tempera- ture a colorless gas is given off from it, and globules of mercury are formed at the same time. The gas, as will be shown later, is oxygen, so that from the red oxide of meroury, which is a chemical compound of mercury and oxygen, we get, by heating, the two elements in the free state. In this case, heat overcomes the chemical attrac- tion which, in the compound, holds the elements together. /Chemical Affinity. — It is evident from what has already been said that there is some power which can hold ele- ments together in a very intimate way, so intimate that we cannot recognize them by ordinary means. We do not know what causes the sulphur and iron to combine, but we do know that they combine. Similarly, we do not know what causes a stone thrown in the air to fall back again, but we know that it falls back. It is true we may say that the cause of the falling of the stone is the attrac- tion of gravitation, but this does not give us any real information, for, if we ask what the attraction of gravita- tion is, we can only answer that it is that which causes all bodies to attract one another. We may also say, and do say, that the cause of chemical combination is chemical affinity. But in so doing we only give a name to something about which we know nothing except the effects it pro- duces. All the chemical changes that are taking place around us may, then, be referred to the operation of chemical affinity. If this power should cease to operate, what would be the result ? Nature would be infinitely less complex than it now is. All complex substances would be resolved into the elements of which they are composed, and, as far as we know, there would be only about seventy different kinds of substances.' All living things would cease to exist, and in their place there would be three in- visible gases, and something very much like charcoal. Mountains would crumble to pieces, and all water would IO COLLEGE CHEMISTRY. disappear, giving two invisible gases. The processes of* life in its many forms would be impossible, as, however subtle that which we call life may be, we cannot imagine it to exist without the existence of certain complex forms of matter; and, as regards the life process of animals and plants, complex chemical changes are constantly taking place within them, and these changes are essential to the continuance of life. Chemical Compounds and Mechanical Mixtures The substances formed by chemical combination of the ele- ments are called chemical compounds. Most substances found in nature are made up of several others. Wood, for example, is very complex, containing a large number of distinct chemical compounds intimately mixed together. Some of these canbe isolated, but it is impossible to isolate them all with the means at present at our command. Most of the rocks met with are also quite complex, and it is difficult to isolate the constituents. If we look at a piece of coarse-grained granite, we see plainly enough that it contains different things mixed together, and if- it is broken up we can pick out pieces of different substances from the mass. If we now examine a piece of each of the different substances thus picked out of the granite, it appears to be homogeneous, that is to say, we cannot recognize the presence of more than one kind of thing in any one piece. If the piece is carefully selected it may be powdered finely in an agate mortar, and some of the powder examined with a microscope without the presence of more than one substance being recognized. It is possible to isolate three substances from granite by simply breaking it up and picking out the pieces of different kinds. It might therefore be concluded that granite consists of three substances. This is true, but it is not the whole truth. For it is possible by proper means to get simpler sub- stances from each of the three already separated. This is, however, a much more difficult process than the separation first accomplished. To effect the separation of each of the three constituents of granite into its elements requires CHEMICAL COMPOUNDS AND CHANGES. n more powerful means. Substances must be brought in contact with them which act upon them, changing their composition, that is to say, act chemically upon them, and high heat must be used to aid the action.' By skilful work it is, however, possible to separate the three components of granite into their elements. From the above it is evident that substances may be united in different ways. They may be so united that it is a simple thing to separate them by mechanical processes. Or they may be so united that it is impossible to separate them by mechanical processes. By a mechanical process is meant any process which does not involve the use of heat, electricity, or chemical change. Thus, the mechan- ical process made use of in the case of granite consisted in picking out the pieces. The separation of the particles of different sizes by means of a sieve is a mechanical process. The separation of two liquids which do not mix with each other is a mechanical process. Complex sub- stances which can be separated into their components by purely mechanical processes are called mechanical mix- tures. Thus granite is a mechanical mixture of three chemical compounds. Similarly, most natural substances are more or less complex mixtures of chemical compounds, or, much more rarely, of elements. Air, for example, is a mechanical mixture consisting mainly of the two elements nitrogen and oxygen. It is not always an easy matter to distinguish between mechanical mixtures and chemical compounds, as there are mixtures which it is extremely difficult to resolve into their components, and there are, on the other hand, chemical compounds which are ex- tremely unstable. Generally, however, the difference is recognized without serious difficulty. Qualitative and Quantitative Study of Chemical Changes. — In general there are two ways in which chem- ical changes may be studied. Substances may be brought together under a variety of conditions and, if action takes place, the properties of the product or products may then be studied and compared with those of the substances 12 COLLEGE CHEMISTRY. brought together. In the early periods of the history of chemistry the study was almost wholly of this kind. This is called qualitative study. But we may go farther than this, and take into consideration the weights or masses o2 the substances we are dealing with. We should then be studying the changes quantitatively. By means of the quantitative method the Jaw of the indestructibility of matter was placed upon a firm basis. By further use of this method other laws of the highest importance to the science of chemistry were soon brought to light. Law of Definite Proportions The fact that sulphur and. iron combine chemically when a mixture of the two is heated has been referred to. The question whether they combine in all proportions is one that can be answered only by a quantitative study of the process. If the process were to be studied for the first time the method of pro- cedure would be this: "We should mix the elements in different proportions and, after the action, we should determine whether any of either of the elements is left in the uncombined state; and, further, by decomposing the product, we should determine whether it always contains the elements in the same proportions. The problem, in this case, is by no means a simple one, but it has been repeatedly worked over with the greatest possible care, and, as the result of the work, the conclusion is justified that the product always contains the elements in exactly the same proportions. Similar work has been done for most other chemical compounds known, and the general conclusion known as the law of definite proportions has been drawn. This law may be stated thus : /"""^ chemical compound always contains the same constit- K_ uents in the same proportion by weight. The truth of this general statement or law has not always been acknowledged by chemists. At the beginning of this century a celebrated discussion on the subject took place between Proust and Berthollet. The discussion led to a great deal of careful work which tended to confirm the law, and since that time it has not been seriously doubted. t LAW OF MULTIPLE PROPORTIONS. 13 About twenty years ago a Belgian chemist, Stas, by a long series of probably the most painstaking and accurate chemical experiments ever performed, showed that in the compounds which he worked with there was no variation in composition that could be detected by the most refined methods of chemistry. In the present state of our knowl- edge it appears that the law of definite proportions is a law in the strictest sense. Law of Multiple Proportions. — It does not require a very extended study of chemical phenomena to show that from the same elements it is possible in many cases to get more than one product. Thus iron and sulphur form three distinct compounds with each other. Tin combines with oxygen in two proportions. The elements potassium, chlorine, and oxygen combine in four different ways, forming four distinct products. Nitrogen and oxygen form five products. In the early part of this century the English chemist, Dalton, by a study of cases like those mentioned was led to the discovery of another great law of chemistry known as the law of multiple proportions. Many substances had been analyzed before his time, and the percentages of the constituents determined with a fair degree of accuracy. He examined, first, two gases, both of which consist of carbon and hydrogen. He determined the percentages of the constituents, and found them to be as follows: Olefiant gas, 85.7 per cent carbon and 14.3 per cent hydrogen. Marsh-gas, 75.0 per cent carbon and 25.0 per cent hydrogen. On comparing these numbers, he found that the ratio of carbon to hydrogen in olefiant gas is as 6 to 1 ; whereas in marsh-gas it is as 3 to 1 or 6 to 2. The mass ot hydrogen, combined with a given mass of carbon, is twice as great in the one case as in the other. There are, further, two compounds of carbon and oxygen, and in analyzing these the following figures were obtained: H COLLEGE CHEMISTRY. Carbon monoxide, 42.86 per cent carbon and 57.14 per cent oxygen. Carbon dioxide, 27.27 per cent carbon and 72.73 per cent oxygen. But 42.86 : 57.14 :: 6 : 8 and 27.27 : 72.73 :: 6 : 16. The mass of oxygen combined with a given mass of car- bon in carbon dioxide is twice as great as the mass of oxy- gen combined with the same mass of carbon in carbon monoxide. These facts and other similar ones led to the discovery of the law of multiple proportions, which may be stated thus : If two elements A and B form several compounds with each other, and we consider any fixed mass of A, then the different masses of B which combine with the fixed mass of A bear a simple ratio to one another. By way of further illustration we may take the three compounds which iron forms with sulphur. In one of these, approximately 7 parts by weight of iron are in com- bination with 4 parts of sulphur; in a second, 7 parts of iron are in combination with 6 parts of sulphur; and in the third, 7 of iron are in combination with 8 of sulphur. The figures 4, 6, and 8 bear a simple ratio to one another which is 2:3:4. The five compounds of nitrogen and oxygen contain 7 parts by weight of nitrogen combined with 8, 16, 24, 32, and 40 parts of oxygen respectively. The figures representing the parts by weight of oxygen combined with 7 parts by weight of nitrogen are in the ratio 1:2:3:4:5. In the compounds formed by the elements chlorine, potassium, and oxygen the proportions by weight are represented in the following table : Chlorine. Potassium. Oxygen, 35.45 39.15 16 35.45 39.15 32 35.45 39.15 48 35.45 39.15 64 It will be observed that the ratio between the chlorine and potassium remains constant, but that the mass of COMBINING WEIGHTS OF THE ELEMENTS. 15 oxygen varies regularly from 16 to 64; the masses bearing to one another the simple ratio 1:2:3:4. The law of multiple proportions like the law of definite proportions is simply a statement in accordance with what has been found true by experiment. Although discovered by Dalton at the beginning of this century and put forward upon what appears now to be only a slight basis of facts, all work since that time has confirmed it/ and it forms to-day one of the corner-stones of the science of chemistry. Combining Weights of the Elements. — A careful study of the figures representing the composition of chemical compounds reveals a remarkable fact regarding the relative quantities of one and the same element which enter into combination with different elements. The proportions by weight in which some of the elements combine chemically with one another are stated in the following table : 1 part Hydrogen combines with 1 it a it tt 35.45 parts Chlorine combine with 39.15 79.96 " Bromine 126.85 tt Iodine 16 a Oxygen 16 a tt 16 tt tt 16 tt tt 65.4 tt Zinc 24.36 tt Magnesium 40' it Calcium 137.4 t* Barium 35.45 parts Chlorine. 79.96 tt Bromine. 126.85 tt Iodine. 39.15 tt Potassium. 39.15 ti tt 39.15 It tt 65.4 t t Zinc. 24.36 ■ ( Magnesium. 40 a Calcium. 137.4 a Barium. 32.06 K Sulphur. 32.06 1 ( tt 32.06 it tt 32.06 tt tt It will be seen that the figures that express the relative weights of chlorine, bromine, and iodine that combine with 1 part by weight of hydrogen also express the relative weights of these elements that combine with 39.15 parts by weight ot potassium. So also the figures which express the relative weights of zinc, magnesium, calcium, and barium that combine with 16 parts by weight of oxygen also express the relative weights of these elements that combine with 32.06 parts by weight of sulphur. Now, an 1 6 COLLEGE CHEMISTRY. examination of all compounds known has shown that hydrogen enters into combination with the other elements in the smallest proportions; it is therefore taken as unity in stating the relative weights of the other elements which enter into combination. The weight of another element that combines with 1* part by weight of hydrogen may be called its combining tueight. Thus, according to the above, the combining weights of chlorine, bromine, and iodine are respectively 35.45, 79.96, and 126.85. Similarly 39.15 is the combining weight of potassium, as it expresses the weight of potassium that combines with the above weights of chlorine, bromine, and iodine. Thus for every element a number can be selected, such that the proportions by weight in which the element enters into combination with others can be conveniently expressed by this number or by a simple multiple of it. These numbers are the com- bining weights. It is not by any means an easy matter to determine which numbers are most convenient for all circumstances; and if the selection is to be determined solely by conve- nience, there may be differences of opinion as to what is most convenient. We shall see a little later that, while the numbers primarily express the combining weights and nothing else, and are based solely upon a study of the composition of chemical compounds, they have come to have a deeper significance which will be explained in due time. Those adopted are called atomic weights because for strong reasons they are believed to express the relative weights or masses of the minute indivisible particles or atoms of which the various kinds of matter are assumed to be made up. The atomic theory, as it is called, will be treated of farther on, and the relation between the theory and the figures called the atomic weights will be discussed at some length. For the present it will be best to use the figures as expressing the combining weights, and as being entirely independent of any speculations regarding the constitution of matter and the existence of atoms. *Seep. 19. SYMBOLS AND ATOMIC WEIGHTS OF THE ELEMENTS. 17 The Elements, their Symbols and Atomic Weights It has already been stated that there are about seventy elementary substances known, but that of these only a small number enter into the composition of common things to any great extent. It has been calculated that the solid crust of the earth is made up approximately as repre- sented in the subjoined table: Calcium 3. 77 7=-— — - and V = 273 + t 1 + 0.00366*" This is called the law of Dalton and Gay-Lussac. On lowering the temperature of a gas below 0° the volume is diminished at the same rate, that is to say it loses ^ part of its volume at 0° for each degree. If this diminution in volume should be regular, it is obvious that at a temperature of — 273° below 0° the volume of a gaseous substance would be reduced to nothing.' This temperature is therefore called absolute zero. The nearest approach to this temperature thus far produced is 16° above the absolute 0°. This was produced by Dewar by boiling liquid hydrogen. It is customary to reduce the observed volume of a gas to the volume which it would have at 0°. The correction is easily made by the aid of the above formula. Thus, if the volume of a gas is found to be 250 cubic centimetres at 15°, and it is required to know what the volume would be if the temperature were reduced to 0°, the calculation is made thus : In this case the observed volume V is 250 cc. ; *, the temperature, is 15°. Substituting those values in the equation V = — or V = 273 + *' 1 + 0.00366*' we have ^ 273 x 250 „ 250 V= -„„ . , „ , or V=: 273 + 15 ' 1 + 0.00366 x 15 ' from which we get 236.99 as the value of V. But the volume of a gas varies also according to the pressure. If the pressure is doubled, the volume is decreased one half ; and if the pressure is decreased one half, the volume is doubled, and so on. In other words, the volume of a gas varies inversely ac- cording to the pressure. Increase the pressure two, three, or four times, and the volume becomes one half, one third, or one fourth, and vice versa. If the gas has the volume V at the pressure P, and at pressure P' the volume V, these values are found to bear to one another the relations expressed in the equation VP = V'P'- This is called Boyle's law. The pressure is usually stated in millimetres, and reference is to the height of a column of mercury which the pressure corresponda MEASUREMENT OF VOLUME OF GASES. 45 to. A gas contained in an open vessel, or in a vessel over mer- cury or water, in which the level of the liquid inside and outside the vessel is the same, is under the pressure of the atmosphere. What that is we learn from the barometer. As this pressure varies, it is necessary to read the barometer whenever a gas is measured, and then to reduce the observed volume to certain conditions which are accepted as standard. If the gas is meas- ured in a tube over mercury or water, and the level of the liquid inside the tube is higher than that outside, the gas is under diminished pressure, the amount of diminution depending on the height of the col- umn of mercury or water in the tube. Thus, if the arrangement is as represented in Fig. 4, the height of the mercury column above the level of the mercury in the trough being 100 millimetres, and the pres- sure of the atmosphere 760 milli- metres of mercury ; then the gas in the tube is plainly not under the full atmospheric pressure, for the atmos- phere is supporting a column of mercury 100 millimetres high, and the pressure actually brought to bear on the gas corresponds to 760 — 100 = 660 mm. of mercury. Suppose that in this case the volume of gas actually measured is 75 cc. Call this V. "What would be the actual volume V under the standard 760 mm. ? "We have seen that VP = V'P' Now, in this case P= 760, V - 75, and P' - 660. Therefore-, 760F=75 x 660, or 7= * 6Q = 65.13. In all cases it is necessary to make a correction similar to this in dealing with the volumes of gases. The correction for tem- perature and that for pressure may be made in one operation, the formula being _ V'P' V = Fie. 4. 760(273 + *)' or 7= 760(1 + 0.003662) ' in which V= the volume of the gas at 0° and 760 mm. pressure ; V = the observed volume ; t = the observed temperature ; P' = the pressure under which the gas is measured. Some of 4 6 COLLEGE CHEMISTRY. B the most important ideas that have been introduced into chem- istry with a view to explaining the regularities observed in the quantities of substances which act upon one another chemically have their origin in observations on the conduct of gases. It is therefore of the highest importance that the student should familiarize himself with the meaning of the expression, " the volume of a gas under standard condi- tions." The presence of water-vapor in a gas also influences its volume, increasing it by the same amount as it would be in- creased if the pressure were diminished by a quantity equal to the pressure of the water-vapor for the given temperature. The formula for making all the correc- tions required in determining the volume of a gas is mV{P'-a) V'(P'-a) 760(273+0 ' - 760(1+0.00366* in which the letters V, V, P', and t have the same significance as in the last formula given, while a is the pressure of water- vapor at t°. The pressure of water-vapor was formerly called " aqueous tension," and this term is still often used. The amount of the pressure for various tem- peratures is to be found from tables. A convenient apparatus for measuring gas volumes, which simplifies the process, is that represented in Fig. 5. It consists of two tubes connected at the base by means of a piece of rubber tubing, and containing water. The tube A is gradu- ated, the other is not. The gas the vol- ume of which is to be measured is brought into the tube A, with the narrow opening at the top, and the other tube is then placed at the side of the one containing the gas, and its height adjusted so that the column of liquid in both tubes is at the same level. Under these circum- stances, obviously, the gas is under the atmospheric pressure for Fiq. 5. DETERMINATION OF AMOUNT OF OXYGEN. 47 which the necessary correction must of course be made. It is also necessary in this case to make the corrections for tempera- ture and the pressure of aqueous vapor. It is, further, some- times convenient when the gas is measured over water to transfer the measuring-tube to a vessel containing enough water to permit the immersion of the tube to a point at which the level of the liquid inside and outside of the tube is the same. In this case the conditions are the same as in the apparatus described in the last paragraph. The arrangement is represented in Fig. 6. Determination of the Amount oe Oxygen liberated when a Known Weight oe Potassium Chlorate is decomposed by heat. Experiment 17. — To determine how much oxygen is given off when a known weight of potassium chlorate is decomposed by heat, proceed as follows : In a small dry hard-glass tube about 10 cm. long and 5 to 7 mm. internal diameter, closed at one end, weigh out on a chemical balance about 0.25 gram dry potassium chlorate, first weighing the tube empty. Introduce just above the potassium chlorate a plug of asbestos which has been ignited; then, by means of a blast-lamp, soften the upper end of the tube, and draw it out so that it has the form shown in Fig. 7. Now weigh the tube again. Let a = weight of tube empty ; I = weight of tube with potassium chlorate ; c = weight of tube with potassium chlorate and plug. Fig. 7. A Fig. 0. Connect at A by means of a short piece of rubber tubing with 48 COLLEGE CHEMISTRY. the measuring-tube Fig. 5 so that the ends of the two tubes are almost in contact with each other, the measuring-tube having been previously filled with water to the zero point, and the top closed by means of the stop-cock. Open the stop-cock, and now heat the potassium chlorate gently at first, and gradually higher until no more gas is given off. After the gas has stood for half an hour to cool it down to the temperature of the air, adjust the two tubes of the measuring apparatus so that the level of the water in both is the same ; read off the volume of gas. At the same time read the barometer and thermometer ; and now make the corrections for pressure and temperature as above di- rected. The weight of a litre or 1000 cc. of oxygen at 0° and 760 mm. pressure is 1.429 grams. Knowing the volume of oxygen obtained, calculate the weight of this volume. Remove the tube containing the product left after the decomposition of the potas- sium chlorate and weigh it. Let d = weight of tube after decomposition of potassium chlorate. Now 6 — a = weight of potassium chlorate used ; c — d~= weight of oxygen given off. This weight should, of course, be the same as that calculated from the measured volume of oxygen, and should further be in agreement with the results calculated by means of the equation KC10 3 = KC1 + 30. Make all the calculations, and see how nearly the results ob- tained agree with what is required by this equation. Should the results not be satisfactory the first time, repeat the work. The more carefully the work is done, the more nearly will the results agree with the equation. „ Experiment 18.— Mix 25 to 30 grams (or about an ounce) of potassium chlorate with an equal weight of manganese dioxide in a mortar. The substances need not be in the form of powder. Heat the mixture in a glass retort, and collect the gas by displace- ment of water in appropriate vessels,— cylinders, bell-glasses, bottles with wide mouths, etc. It will also be well to collect some in a gasometer, such as is commonly found in chemical laboratories, the essential features of which are represented in Fig. 8. It is made either of metal or of glass. The opening at d can be closed by means of a screw cap. In order to fill it DETERMINATION OF AMOUNT OF OXYGEN. 49 with water, open the stop-cocks and pour the water into the upper part of the vessel after hav- ing screwed on the cap d. When it is full, water will flow out of the small tube e. Now close all the stop-cocks, and remove the cap d. The water will stay in the vessel for the same reason that it will stay in a cylinder in- verted with its mouth below water. To fill the gasometer with gas, put it over a tub or sink, and introduce the tube from which gas is issuing into the opening at d. The gas will rise and displace the water, which will flow out at d. When full, put the cap on. To get the gas out of the gasometer, attach a rubber tube to e, pour water into the upper part of the gasometer, open the stop-cock a and that at e, when the gas will flow out, and the current can be regulated by means of the stop-cock at e. The arrangement of the retort is shown in Fig. 9. Fig. 8. Fig. 9. 5° COLLEGE CHEMISTRY. Physical Properties of Oxygen. \ Experiment 19.— Inhale a little of the gas from one of the bottles. Has it any taste ? odor ? color ? Chemical Properties of Oxygen. Experiment 20.— Turn three of the bottles containing oxygen with the mouth upward, leaving them covered with glass plates. Into one introduce some sulphur in a so- called deflagrating-spoon, which is a small cup of iron or brass attached to a stout wire which passes through a metal plate, usually of tin (see Fig. 10). In another put a little charcoal (carbon), and in a third a piece of phosphorus* about the size of a pea. Let them stand quietly and notice what changes, if any, take place. Sulphur, carbon, and phos- phorus are elements, and oxygen is an element. It will be noticed that the sul- phur and the carbon remain unchanged, while some change is taking place in the vessel containing the phosphorus, as is shown by the appearance of white fumes. After some time the phosphorus will disappear entirely, the fumes will also disappear, and there will be nothing to show us what has become of the phosphorus. If the temperature of the room is rather high, it may happen that the phosphorus takes fire. If it should, it will burn with an intensely bright light. After the burning has stopped, the vessel will be filled with white fumes, but these will quickly disappear, and the vessel will apparently be empty. What do these experiments prove with reference to the action of oxygen on sulphur, carbon, and phos- phorus at the ordinary temperature ? * Phosphorus should be handled with great care. It is always kept under water, usually in the form of sticks. If a small piece is wanted, take out a stick with a pair of forceps, and put it under water in an evaporating-dish. While it is under the water, cut off a piece of the size wanted. Take this out by means of a pair of forceps, lay it for a moment on a piece of filter-paper, which will absorb most of the water, then quickly put it in ths spoon. Fio. 10. CHEMICAL PROPERTIES OF OXYGEN. 51 \ Experiment 21. — In a deflagrating-spoon set fire to a little sulphur and let it burn in the air. Notice whether it burns with ease or with difficulty. Notice the odor of the fumes which are given off. Now set fire to another small portion and intro- duce it in a spoon into one of the vessels containing oxygen. It will be seen that the sulphur burns much more readily in the oxy- gen than in the air. Notice the odor of the fumes given off. Is it the same as that noticed when the burning takes place in the air ? " v Experiment 22. — Perform similar experiments with charcoal. ,, Experiment 23. — Burn a piece of phosphorus not larger than a pea in the air and in oxygen. In the latter case the light emitted from the burning phosphorus is so intense that it is pain- ful to some eyes. It is best to be cautious. The phenomenon is an extremely brilliant one. The walls of the vessel in which the burning takes place become covered with a white substance which afterwards gradually disappears. What differences do you notice between the burning in the air and in oxygen ? In the experiments is there any sulphur,'or car- bon, or phosphorus left behind ? Do the experiments furnish any evidence that oxygen takes part in the action ? or that oxygen is used up ? , Experiment 24. — Straighten a steel watch-spring * and fasten it in a piece of metal, such as is used for fixing a deflagrating- spoon in an upright position ; wind a little thread around the lower end, and dip it in melted sulphur. Set fire to this and in- sert it into a vessel containing oxygen. For a moment the sul- phur will burn as in Experiment 21 ; but soon the steel begins to burn brilliantly, and the burning continues as long as there is oxygen left in the vessel. Notice that in this case there is no flame, but instead very hot. particles are given off from the burn- ing iron. The phenomenon is of great beauty, especially if ob- served in a dark room. The walls of the vessel become covered with a dark reddish-brown substance, some of which will also be found at the bottom in larger pieces. * Old watch-springs can generally be had of any watch maker or mender for the asking. A spring can be straightened by unrolling it, attaching a weight, and suspending the weight by the spring. The spring is then heated up and down to redness with the flame of a Bunsen burner. 52 COLLEGE CHEMISTRY. Oxygen is used up in Combustion. Experiment 25.— Is the odor of the contents of the bottle in which the sulphur was burned the same as before the experiment 2 Introduce a stick with a small flame on it successively into the vessels used in burning sulphur, carbon, phosphorus, and iron. Is oxygen present or not? What evidence have you on this point ? , Experiment 26. — Fill a tube say 30 to 40 cm. (12 to 15 inches) long, and 2£ to 3 cm. (1 to 1J inches) wide, with oxygen, and ar- range it in a vessel over water, as shown in Fig. 11. Now fasten a small stick of phosphorus to the end of a wire and push it into the tube so that about J- to J inch of the phosphorus is above the water and exposed to the oxygen. At first no action will take place, but after a time white fumes will be seen to rise from the phosphorus, and the phosphorus will begin to melt. This action will be accompanied by a diminution of the volume of the oxygen, as will be shown by the rise of the water. When the water has risen so as to cover the phosphorus, shove the stick up so that it is again just above the surface of the water. Some of the oxygen will again be used up. By working carefully, and re- peating this process as many times as may be necessary, the oxygen can all be used up with- out the active burning of the phosphorus. Usually, however, before the action is com- pleted, the temperature of the phosphorus be- comes so high that it takes Are, when there is a flash of light in the tube and a sudden rise of the water, showing that the gas is suddenly used up. Experiment 27. — Burn a steel watch-spring as directed in Experiment 24, with the difference that the spring is passed air-tight through a cork which is fitted tightly into the neck of the bell-jar (Fig. 12). As the spring burns, the water will rise from the vessel in which the bell-jar is standing, and it is necessary to pour water into this vessel. When the spring has burned near to the cork shove it through so that the burning may continue. If the experiment is properly performed the bell-jar will be nearly full of water at the end. What does this prove? PRODUCTS OF COMBUSTION. 53 The Products oe Combustion weigh more than the Substance Burned. Experiment 28. — Weigh off about a gram of magnesium ribbon in a porcelain crucible. Heat over a Bunseu burner until the magnesium has turned to a white sub- stance (magnesium oxide). After cool- ing, weigh again. Perform the same ex- periment with zinc, tin, and lead. What conclusion are you justified in drawing ? Experiment 29. — Over each pan of a large and rather sensitive balance suspend a glass tube filled with pieces of solid caustic soda. A balance that will answer the purpose very well can be made of wood with metal bearings. It may con- veniently be about 2£ feet high, with a delicate beam about 3 feet long. The best tubes for the caustic soda are Argand lamp-chimneys, around the bottom of which is tied a piece of wire-gauze to pre- vent the caustic soda from falling out. On one pan of the balance place a candle directly under one of the caustic-soda tubes, so adjusted that the flame shall be not more than 2-J to 3 inches below the bottom of the tube. By means of weights placed on the other pan establish equilibrium. Now light the candle. Slowly, as it burns, the pan upon which it is placed will sink, showing that the products of combustion which are partly ab- sorbed by the caustic soda are heavier than the candle was. While this is by no means an accurate experiment, it is a very striking one, and proves beyond question that in the process of combustion matter is taken up by the burning body. Fig. 12. OHAPTEE III. A STUDY OF THE ELEMENT HYDROGEN. Historical. — Hydrogen was discovered as a distinct sub- stance by Cavendish in 1766, although it had been observed as an inflammable gas before that time. Occurrence. — It occurs to some extent in the free con- dition, and issues from the earth in small quantity in some localities, as, for example, in the salt-mines at Stassfurt, Germany. It has also been shown to occur in enormous quantities in the atmosphere of the sun. On the earth it occurs chiefly, however, in combination in water, of which it forms 11.19 per cent. It occurs also in most substances of animal and vegetable origin, such as the various kinds of wood and fruits, and the tissues of all animals. In these products of life it is contained in com- bination with carbon and oxygen or with carbon, oxygen, and nitrogen. Preparation. — The simplest way, theoretically, to pre- pare hydrogen is by the decomposition of water by the electric current. But this method is less convenient and more expensive than other methods which are available, and it is therefore used only under special circumstances. It is particularly well adapted to the preparation of small quantities of pure hydrogen. Some elements, as, for example, sodium and potassium, when brought in contact with water at the ordinary tem- perature decompose it and set hydrogen free. When a small piece of potassium is thrown upon water, a flame is observed at once. When sodium is used, it is seen to form a small ball which moves about on the surface of the water 54 PREPARATION OF HYDROGEN. 55 with a hissing sound, but under ordinary circumstances no flame is observed. By applying a flame to the ball something takes fire and burns. By filling a good-sized test-tube with water and inverting it in a larger vessel and bringing a small piece of sodium wrapped in a piece of filter-paper below the mouth of the tube, the sodium will rise to the top of the tube when released, and it will then be seen that a gas is evolved which gradually depresses the water in the tube. A similar experiment with potassium gives a similar result. The gas given off in each case is hydrogen. By evaporating off the water left in the vessel there will be found in each case a white substance of marked chemical properties. That formed with the potas- sium is known as potassium hydroxide, or caustic potash, and has the composition represented by the symbol KOH; that formed with the sodium is known as sodium hy- droxide, or caustic soda, and is represented by the symbol NaOH. The reactions between potassium and sodium and water are represented by the equations K + H 2 = KOH + H; and Na + H 2 = NaOH + H. Half the hydrogen of the water which is decomposed is replaced by the potassium or the sodium, as the case may be. These reactions are partly described by saying that the potassium or sodium is substituted for half the hydrogen in the water, and the act is called substitution. This is a very common kind of chemical action, and we shall constantly meet with it in the course of our study. The reaction is one of double decomposition or metathesis, two substances acting upon each other to form two others. Some substances, such as iron, which decompose water slowly at the ordinary temperature do so readily at higher temperatures. At ordinary temperatures iron decomposes water, as is seen in the formation of a coating upon it when left in contact with water and the air. At higher temperatures when the iron is red-hot it decomposes water very readily, and hydrogen can be made in quantity by 56 COLLEGE CHEMISTRY. this means. In the laboratory the iron may be heated in a gun-barrel or in a porcelain tube. When steam is passed over it the decomposition represented in this equation takes place : 3Fe + 4H 2 = Fe 3 4 + 8H. The iron combines with the oxygen and liberates the hydrogen. Carbon, in the form of charcoal or coal, may be used in a similar way to effect the decomposition of water and the liberation of hydrogen. At a high heat the reaction takes place mainly as represented thus : C + H a O = CO + 2H. A mixture of two gases, carbon monoxide and hydrogen, is thus formed. This mixture is the essential part of the gas which has of late years come into such extensive use under the name "water-gas." This is formed by passing steam over highly-heated anthracite coal or coke. By far the most convenient method for making hydrogen consists in treating a metal with an acid. Among the metals best adapted to the purpose are zinc and iron, and indeed zinc is almost exclusively used. As will be seen later, acids are substances that contain hydrogen, and are characterized by the property that they give up this hydrogen easily and take up other elements in the place of it. Among the common acids found in every laboratory are hydrochloric acid, sulphuric acid, and nitric- acid. The chemistry of these compounds will be treated of in due time; but, as we shall be obliged to use them before they are taken up systematically, a few words in regard to them are desirable in this place. Hydrochloric acid is a compound of hydrogen and chlorine. It is a gas which dissolves easily in water. It is this solution that is used in the laboratory. It is manufactured in enormous quantities in connection with the manufacture of soda or sodium carbonate. Its chemi- PHYSICAL PROPERTIES OF HYDROGEN. 57 cal symbol is HC1. In commerce it is not uncommonly called "muriatic acid." Sulphuric acid is a compound of sulphur, oxygen, and hydrogen in the proportions represented by the formula H 2 S0 4 . It is an oily liquid and is frequently called "oil of vitriol." It is manufactured in very large quantities, as it plays an important part in many of the most impor- tant chemical industries. Nitric acid is a compound containing nitrogen, oxygen, and hydrogen in the proportions represented by the formula HN0 3 . It is a colorless liquid, though, as we get it, it is commonly colored straw-yellow. When a metal, such as zinc, is brought in contact with hydrochloric or sulphuric acid, an evolution nf hydrogen takes place at once. The reactions are as represented in these equations : Zn + 2HC1 = ZnCl 2 + 2H; Zinc Chloride. Zn + H 2 S0 4 = ZnSO, + 2H. Zinc Sulphate. Each combining weight of zinc liberates and replaces two combining weights of hydrogen. The action between iron and these two acids is of the same character: Fe + 2H01 = FeCl 2 + 2H; FeiTous Chloride. Fe + H 2 S0 4 = FeS0 4 + 2H. Ferrous Sulphate. The hydrogen obtained from acids by the action of metals is not pure, but it can be purified by treatment with appropriate substances. That obtained by the decomposition of water by the electric current is pure. Physical Properties. — Hydrogen is a colorless, inodor- ous, tasteless gas. That made by the action of acids on zinc or iron has a somewhat disagreeable odor which is due to the presence of other gases in small quantity. It is not poisonous, and may therefore be inhaled with impunity. We could not, however, live in an atmosphere of hydrogen, 58 COLLEGE CHEMISTRY. as we need oxygen. It is the lightest known substance. Its specific gravity in terms of the air standard is 0.0696. A litre under 760 mm. pressure and at 0° weighs 0.08995 gram. Under Oxygen it was stated that in chemistry hydrogen is commonly taken as the standard of specific gravity, and that, hydrogen being unity, the specific gravity of oxygen is 15.89. The gas is only slightly solu- ble in water. 100 volumes of water take up 1.93 volumes of hydrogen. The fact that hydrogen is lighter than the air is shown by opening a vessel which contains it and turning the mouth of the vessel upward. The gas escapes and in a very short time no evidence of its pres- ence can be obtained. Light vessels, as, for example, soap-bubbles or collodion-balloons, filled with the gas rise in the air, and it is used for the purpose of filling large balloons. At a very low temperature and high pressure, hydrogen can be converted into a liquid that boils at — 252°. It cannot be liquefied at any temperature above — 242°, no matter what pressure it may be subjected to. Hydrogen passes readily through porous substances, or it diffuses rapidly. This can easily be demonstrated in the case of porous earthenware and paper. It also passes readily through some metals, as iron and platinum, when heated to redness. There is a direct relation between the specific gravity of gases and the rate at which they diffuse. The lower the specific gravity the more rapid the diffusion. The law governing these phenomena is : The rate of diffusion of gases is approximately inversely proportional to the square roots of their specific gravities. The specific gravity of hydrogen being 1 and that of oxygen nearly 16 (15.89), the rate of diffusion of oxygen is approximately ^ that of hydrogen. If hydrogen is on one side of a porous wall, and oxygen on the other, the hydrogen will pass through the wall so much more rapidly than the oxygen that there will be an accumulation of hydrogen on one side of the wall, and if the vessel were CHEMICAL PROPERTIES OF HYDROGEN. 5^ closed there would be increased pressure on that side. The ready passage of gases through porous walls is a matter of great importance in connection with the ventila- tion of dwellings. Most of the materials used in building are porous and permit the passage of gases through them in both directions, and change of air is secured in this way to some extent. Chemical Properties. — Under ordinary circumstances hydrogen is not a particularly active element. It does not unite with oxygen gas at ordinary temperatures, but, like other combustible substances, it must be heated to the kindling temperature before it will burn. If a flame is applied to it, it takes fire at once. The flame is colorless or slightly blue. Generally the flame is somewhat colored in consequence of the presence of foreign substances; but that it is colorless when the gas is burned alone can be shown by burning it as it issues from a platinum tube which is itself not chemically acted upon by the heat. Although the flame is not luminous it is intensely hot, as can be seen by inserting into it a coil of platinum wire, which will at once become red-hot and emit light accord- ingly. The burning of hydrogen in the air, like the burning of other combustible substances in the air, consists in a union of the gas with oxygen. This has been shown to be true by most elaborate experiments on the combustion of hydrogen in oxygen and in the air. On the other hand, substances which burn in the air are extinguished when put in a vessel containing hydrogen. This is equivalent to saying that a substance which is uniting with oxygen does not continue to unite with oxygen when put in an atmosphere of hydrogen, and does not combine with hydrogen. The fact is expressed by saying that hydrogen does not support combustion. This can be shown by holding a vessel filled with hydrogen with the mouth downward, and inserting into it a lighted taper supported on a wire. The gas takes fire at the mouth of the vessel, but the taper is extinguished. 60 COLLEGE CHEMISTRY. Ordinarily we say that hydrogen burns in oxygen, but, as the act consists in the union of the two gases, it would seem probable that oxygen will burn in an atmosphere of hydrogen. This can be shown to be true by a proper arrangement of apparatus. If we were surrounded by an atmosphere of hydrogen we should speak of oxygen as a combustible gas in the same way that we now speak of hydrogen as a combustible gas. It can easily be shown that, when hydrogen is burned either in oxygen or air, water is formed. The simplest way to show this is by holding a glass plate or some other incombustible object a short distance above a flame of hydrogen. It will be seen that drops of water are con- densed upon it. Hydrogen combines with many other elements besides oxygen, and forms some of the most important and inter- esting compounds, such as hydrochloric acid, HC1; sul- phuretted hydrogen or hydrogen sulphide, H 2 S ; ammonia, NH 3 ; marsh-gas, 0H 4 ; and all the acids. On account of its power of combining with oxygen it is used very exten- sively in the laboratory for the purpose of extracting oxygen from compounds containing it. Thus, when hydrogen is passed over heated copper oxide, OuO, it combines with the oxygen to form water, and the copper is left in the free or uncombined state. The reaction is represented thus : CuO + 2H = H 2 + Cu. A similar reaction takes place when hydrogen is passed over highly-heated oxide of iron, Fe.,0,,: Fe 2 3 + 6H = 2Pe + 3H 2 0. The removal of oxygen from a compound is called reduction. Eeduction is therefore plainly the opposite of oxidation. Any substance that has the power to abstract oxygen is spoken of as a reducing agent, just as any sub- stance that has the power to add oxygen to a substance, COMPARISON OF OXYGEN AND HYDROGEN. 6 1 or to decompose it by the action of oxygen, is called an oxidizing agent. A number of metals have the power to absorb a large quantity of hydrogen when they are heated in a current of the gas. This phenomenon is shown most strikingly by palladium, which under the most favorable conditions takes up something more than 935 times its own volume of hydrogen. The gas is given up at an elevated tempera- ture in a vacuum. When it absorbs hydrogen, palladium undergoes marked changes in properties. Its volume is increased, and its magnetic and electric properties are also changed. It was suggested by Graham that the hydrogen held in combination by the palladium is something quite different from ordinary hydrogen, and that it must have some properties like those of the so-called metals. He therefore called the combined hydrogen hydrogenium. Later studies, especially of liquid hydrogen, have shown that there is little in common between hydrogen and the metals. Comparison of Oxygen and Hydrogen. — Hydrogen and oxygen are different kinds of matter, just as heat and electricity are different kinds of energy. Heat can be converted into electrical energy, and electrical energy into heat, but one element cannot by any means known to us be converted into another. They are apparently entirely independent of each other. The question will therefore suggest itself whether, in spite of their apparent inde. pendence, there is not some relation between the different elements which reveals itself by similarity in properties. It will be found that the elements can be separated into groups or families according to their properties. There are some elements, for example, which in their chemical conduct resemble oxygen markedly. These elements con- stitute the oxygen family. So far as hydrogen is con- cerned, however, 'it stands by itself. There is no other element which conducts itself like it. If we compare it with oxygen, we find very few facts that indicate any analogy between the two elements. In their physical properties they are, to be sure, similar. Both are color- 62 COLLEGE CHEMISTRY. less, inodorous, tasteless gases. On the other hand, oxygen combines readily with a large number of substances with which hydrogen does not combine. Oxygen, as we have seen, combines easily with carbon, sulphur, phosphorus, and iron. It is a difficult matter to get any of these ele- ments to combine directly with hydrogen. Further than this, substances which combine readily with hydrogen do not combine readily with oxygen. The two elements ex- hibit opposite chemical properties. What one can do the other cannot do. This oppositeness of properties is favor- able to combination; for not only do hydrogen and oxygen combine with great ease under proper conditions, but, aa we shall see later, it is a general rule that elements of like properties do not readily combine with one another, while elements of unlike properties do readily combine. EXPERIMENTS. Preparation of Hydrogen - . Experiment 30. — Eepeat Experiment 3 and examine the gases. Experiment 31. — Throw a small piece of sodium * on water. While it is floating on the surface apply a lighted match to it. A yellow flame will appear. This is burning hydrogen, the flame being colored yellow by the presence of the sodium, some of which also burns. Make the same experiment with potassium. The flame appears in this case without the aid of the match. It has a violet color, which is due to the burning of some of the potas- sium. The gas given off in these experiments is either burned at once or escapes into the air. In the case of the potassium it takes fire at once, because the action takes place rapidly and the heat evolved is sufficient to set fire to it ; in the case of the sodium, however, the action takes place more slowly, and the temperature does not getjhigh enough to set fire to the gas. In order to collect it unburned, it is only necessary to allow the decomposition to take place so that the gas will rise in an inverted vessel filled with * The metals sodium and potassium are kept under oil. When a small piece is wanted, take out one of the larger pieces from the bottle, roughly wipe off the oil with filter-paper, and cut off a piece the size needed. It is not advisable to use a piece larger than a small pea. PREPARATION OF HYDROGEN. 63 water. For this purpose fill a good-sized test-tube with water and invert it in a vessel of water. Cut off a piece of sodium not larger than a pea, wrap it in a layer or two of filter-paper, and with the fingers or a pair of curved forceps bring it quickly below the mouth of the test-tube and let go of it. It will rise to the top, the de- composition of the water will take place quietly, and the gas formed, being unable to escape, will remain in the tube. By repeating this operation in the same tube a second portion of gas can be made, and so on until the vessel is full. Examine the gas and see whether it acts like the hydrogen obtained from water by means of the electric current. What evi- dence have you that they are the same ? Is this evidence sufficient to prove the identity of the two ? The metals sodium and potassium disappear in these experi- ments, and we get hydrogen. "What becomes of the metals ? and what is the source of the hydrogen ? If after the action has stopped the water is examined, it will be found to contain something in solution. It now has a peculiar taste, which we call alkaline ; it feels slightly soapy to the touch ; it changes certain vegetable colors. If the water is evaporated off, a white substance remains behind, which is plainly neither sodium nor potassium. In solid form or in very concentrated solution it acts very strongly on animal and vegetable substances, disintegrating many of them. On account of this action it is known as caustic soda, or, in the case of potassium, as caustic potash. Fio. 13. Experiment 32. — Certain metals that do not decompose water at ordinary temperatures, or decompose it slowly, decompose it easily at elevated temperatures. This is true of iron. If steam is passed through a tube containing pieces of iron heated 6 4 COLLEGE CHEMISTRY. to redness, decomposition of the water takes place, and the oxy- gen is retained by the iron, which enters into combination with it, while the hydrogen is liberated. In this experiment an iron tube with an internal diameter of from 20 to 25 mm. (about an inch) and a gas-furnace are desirable, though a hard-glass tube and a charcoal-furnace will answer. The arrangement of the apparatus is shown in Fig. 13. Experiment 33. — In a cylinder or test-tube put some small pieces of zinc, and pour upon it some ordinary hydrochloric acid. If the action is brisk, after it has continued for a minute or two apply a lighted match to the mouth of the vessel. The gas will take Are and burn. If sulphuric "acid diluted with five or six times its volume of water * is used instead of hydrochloric acid, Fig. 14. Fig. 15. the same result will be reached. The gas evolved is hydrogen. For the purpose of collecting the gas the operation is best per- formed in a wide-mouthed bottle, in which is fitted a cork with two holes (see Fig. 14), or in a bottle with two necks called a * If it is desired to dilute ordinary concentrated sulphuric acid with water, the acid should be poured slowly into the water while the mixture is constantly stirred. If the water is poured into the acid, the heat evolved at the places where the two come in contact may be so great as to convert the water into steam and cause the strong acid to spatter. * EXPERIMENTS WITH HYDROGEN. 65 "Woulff's flask (see Fig. 15). Through one of the holes a funnel- tube passes, and through the other a glass tube bent in a con- venient form. The zinc used is granulated. It is prepared by melting it in a ladle, and pouring the molten metal from an elevation of four or five feet into water. The advantage of this form is that it pre- sents a large surface to the action of the acids. A handful of this zinc is introduced into the bottle, and enough of a cooled mixture of sulphuric acid and water (1 volume concentrated acid to 6 volumes water) poured upon it to cover it. Usually a brisk evolution of gas takes place at once. "Wait for two or three minutes, and then collect some of the gas by displacement of water. "When the action becomes slow, add more of the dilute acid. It will be well to fill several cylinders and bottles with the gas, and also a gasometer, from which it can be taken as it is needed for experiments. Something besides Hydrogen is formed. Experiment 34.— After the action is over pour the contents of the flask through a filter into an evaporating-dish, and boil off the greater part of the water, so that, on cooling, the substance contained in solution will be deposited. If the operation is car- ried on properly, the substance will be deposited in regular forms called crystals. It is zinc sulphate, ZnSO< , formed by the sub- stitution of zinc for the hydrogen of the sulphuric acid. Problems. — How much zinc would it take to give 200 litres of hydrogen ? How much zinc sulphate would be formed ? How much hydrogen would be formed by the action of 50 grams of zinc on sulphuric acid ? How much sulphuric acid would be used up ? Determination oe the Amount of Hydrogen eyolved when a Known Weight of Zinc is dissolved in Sulphuric Acid. Experiment 35. — This determination can be made by means of an apparatus such as represented in Fig. 16. The bent tube lead- ing from the flask A is drawn out at B, and a plug of glass-wool introduced below the constriction. The other parts of the appa- ratus need no description. The flask should have a capacity of 66 COLLEGE CHEMISTRY. about 40 to 50 cc. ; and the measuring tube C should have a capacity of about 100 cc, and be graduated to ^ cc. " The experiment is conducted in the following manner : D is filled with distilled water ; a piece of zinc weighing from 0.150 to 0.200 gram is placed in the flask ; the pinch-cock E is then opened, and the whole apparatus thus filled with water. The apparatus is now examined in order to ascertain if gas-bubbles are lodged under the stopper F or in the glass-wool. If so, they can usually be dislodged without difficulty. If they persist, a few moments' boiling of the water in the flask will effect their complete re- moval. . . . The eudiometer is now placed over the outlet of the delivery-tube, and the greater portion of the water'remaining in D allowed to flow through the apparatus. Sulphuric acid of the concentration ordinarily employed in the laboratory (1 of HjSC^ to 4 of H,0) is poured into the reservoir D until it is nearly full. The pinch-cock E is then opened, and the water which fills the Fio. 16. apparatus is displaced by sulphuric acid. The action of the acid upon the metal may be facilitated by heat or by adding some platinum scraps*:' When the action is over, the contents of" the flask are swept through the delivery-tube by again opening the pinch-cock E. Finally, the measuring-tube is transferred to a cylinder of water, and the volume of the gas read and corrected in the usual manner. If hydrochloric instead of sulphuric acid has been used, which would be the case when the metal employed is aluminium, a little caustic soda should be added to the water in the cylinder to which the eudiometer is transferred." EXPERIMENTS WITH HYDROGEN. 67 A litre of hydrogen at 0° and 760 mm. weighs 0.0895 gram. How much does the hydrogen obtained in the experiment weigh ? How much ought to have been obtained ? How many cubic cen- timetres of hydrogen ought to have been obtained ? Try the same experiment, using tin and hydrochloric acid. The action takes place as represented in the equation Sn + 2HC1 = SnCl 2 + H„. It would be well, further, to try the experiment also with iron and sulphuric acid, and with aluminium and hydrochloric acid, and to calculate from the results the relation between the weights of the four metals required to give equal volumes of hydrogen, and the volumes of hydrogen given by, say, a gram of each metal. The action between iron and sulphuric acid takes place according to the equation Fe + H 2 S0 4 = FeS0 4 + H„. That between aluminium and hydrochloric acid is represented by this equation : Al + 3HC1 = Aid, + 3H. Fig. 17. 68 COLLEGE CHEMISTRY. Hydrogen is Purified by passing through a Solu- tion of Potassium Permanganate. Experiment 36. — Pass some of the gas, made by the action of zinc on sulphuric acid, through a wash-cylinder containing a solu- tion of potassium permanganate ; collect some of it, and notice whether it has an odor. The apparatus should be arranged as shown in Fig. 17. The solution of potassium permanganate is, of course, contained in the small cylinder A, and the tubes so arranged that the gas bubbles through it. Has the gas any odor or taste or color ? Experiment 37. — Place a vessel containing hydrogen with the mouth upward and uncovered. In a short time examine the gas contained in the vessel, and see whether it is hydrogen. What does this experiment prove with reference to the weight of hy- drogen as compared with that of the air ? Experiment 38. — Gradually bring a vessel containing hydrogen with its mouth upward below an inverted vessel containing air, in the way shown in Fig. 18. After the vessel which contained the hydrogen has been brought in the upright position beneath the other, examine the gas in each vessel. Which one contains the fig. 18. hydrogen ? Experiment 39. — Soap-bubbles filled with hydrogen rise in the air. This experiment is best performed by connecting an ordi- nary clay pipe by means of a piece of rubber tubing with the delivery-tube of a gasometer filled with hydrogen. Small bal- loons of collodion are also made for the purpose of showing the lightness of hydrogen. Hydrogen passes readily through Porous Vessels. Diffusion. Experiment 40. — Arrange an apparatus as shown in Fig. 19a. It consists of a porous earthenware cup, such as is used in gal- vanic batteries, fitted with a perforated cork connected with a glass tube 2 to 3 feet long. The cork must fit air-tight into the mouth of the cup, as well as the tube into the cork. This may be secured by shoving the cork into the cup until its outer sur- EXPERIMENTS WITH HYDROGEN. 69 Fig. 196. face is even with the edge of the cup, and then covering it cave- fully with sealing-wax. Put the lower end of the glass tube through a cork into one neck of a Woulff's bottle contain- ing some water colored with litmus or indigo, so that the end of the tube is above the surface of the water. Through the other neck of the bottle pass a tube slightly bent out- ward and drawn out at the end to a fine opening. Thi3 tube must also be fitted to the bottle by an a air-tight cork, and its lower end must be below the surface of the liquid. Now bring a bell-jar (Fig. 196) containing dry hydrogen over the porous cup, when the liquid will be seen to rise in the short, bent tube that dips below the liquid, and be forced out of it, sometimes with considerable velocity. Withdraw the bell- jar, and bubbles will rise rap- idly from the bottom of the tube which dips under the water, thus showing that air is entering the bottle. This is due to the diffusion of the hydrogen from the porous cup into the air. Explain all that you have seen. Fig. 19a. Fig. 20. Chemical Properties oe Hydrogen. Experiment 41. — If there is no small platinum tube available, roll up a small piece of platinum-foil and melt it into the end of a glass tube, as shown in Fig. 20. Connect the burner thus made with the gasometer containing hydrogen, and after the gas has been allowed to issue from it for a moment, set fire to it. In a short time it will be seen that the flame is practically colorless, and gives no light, That it is hot can be readily shown by hold- ing a piece of platinum wire or a piece of some other metal in it. 7o COLLEGE CHEMISTRY. Experiment 42.— Into the flame of burning hydrogen intro- duce a small coil of platinum wire. What change is observed ? Introduce also a piece of magnesium ribbon. Explain the difference between the two cases. What becomes of the magnesium ? of the platinum ? Experiment 43.— Hold a cylinder filled with hydrogen with the mouth downward. Insert into it a lighted taper held on a bent wire, as shown in Eig. 21. The gas takes fire at the mouth of the vessel, but the taper is extinguished. On withdrawing the taper and holding the wick for a moment in the burn- ing hydrogen, it will take fire, but on putting it back in the hydrogen it will again be extin- guished. Other burning substances should be tried in a similar way. What conclusions Fig. 81. are justified by the last two experiments ? Product formed when Hydrogen is burned. Experiment 44. — Hold a clean, dry glass plate a few inches above a hydrogen flame. What do you observe ? Remove what is deposited upon the plate, and hold the plate again over the flame. Repeat this a number of times. What does the substance deposited upon the plate suggest ? Can you positively say what it is? Reduction. Experiment 45. — Arrange an apparatus as shown in Fig. 22. The flask A contains zinc and dilute sulphuric acid ; the cylinder B a solution of potassium permanganate ; the tube G granulated calcium chloride ; and the cylinder D concentrated sulphuric acid. The object of the potassium permanganate is to purify the hydrogen ; the object of the concentrated sulphuric acid and cal- cium chloride is to remove moisture from the gas. In the tube E put a few pieces of the black oxide of copper, or cupric oxide, CnO. After hydrogen has been passing long enough to drive all the air out of the apparatus (about two or three minutes if there is a brisk evolution) heat the oxide of copper by means of a flame applied to the tube. What change in color takes place ? Try the EXPERIMENTS WITH HYDROGEN. V- action of nitric acid on the substance before the action and after, and note whether there is any difference. What appears in G ? Explain what you have seen. Fig. 22. Experiment 46.— Try the experiment just described, using ferric oxide, or oxide of iron, Fe 3 3 , instead of cupric oxide. What is the common feature in the two reactions ? CHAPTER IY. STUDY OF THE ACTION OF HYDROGEN ON OXYGEN. Burning of Hydrogen. — Attention has already been called to the fact that, when hydrogen burns, water is formed. One of the first questions to be answered is, What relation exists between the weights of the hydrogen burned, the oxygen used up in the burning process, and the water formed ? But to weigh gases accurately and to collect small quantities of water and weigh it are by no means simple operations, and a great deal of work has been done upon the problem under consideration. Method of Dumas. — The first accurate experiments on the combustion of hydrogen are those of Dumas. The method employed by this chemist was the following: He passed carefully-purified hydrogen over heated copper oxide and collected the water formed. The reaction in- volved is that represented by the equation CuO + 2H = H 2 + Cu. I The weight of the oxygen that entered into combination with hydrogen was obtained by weighing the vessel con- taining the copper oxide before and after the experiment. The loss in weight represented the weight of the oxygen abstracted. The water was collected by passing the vapor into an empty glass vessel in which most of it was ' con- densed, and then through tubes containing molten caustic potash and phosphorus pentoxide, substances which have a marked power to absorb water and hold it in combina- tion. In nineteen experiments he obtayied a total amount of water weighing 945.439 grams, and the total amount 72 MORLEY'S METHOD— EUDIOMETRIC METHOD. 73 of oxygen used in the formation of this amount of water was 840.161 grams. According to these results the ratio between hydrogen and oxygen in water expressed in per- centages is : Oxygen 88.864 Hydrogen 11.136 100.000 Morley's Method. — Morley proceeded as follows: He weighed comparatively large volumes of oxygen in the form of gas; he then weighed hydrogen after it had been absorbed by palladium. The hydrogen being expelled from the palladium, it was brought together with the oxygen, with which it was caused to combine by means of electric sparks. The water thus formed was carefully collected in the same apparatus and weighed. Twelve experiments gave results that agreed closely with one another. The mean of these shows that hydrogen and oxgen combine in the ratio of 2 parts of hydrogen to 15.8792 parts of oxygen. Dumas' result expressed in the same way is 2: 15.961. Eudiometric Method. — "When hydrogen and oxygen are mixed at ordinary temperatures no chemical change takes place, and the two gases may be left in contact with each other indefinitely without chemical action. If, however, a spark is brought in contact with the mixture, violent action takes place, accompanied by a flame and explosion. The action consists in the sudden combination of the two gases to form water. It may be illustrated in a number of ways — most simply by filling soap-bubbles with the mix- ture of the two gases and applying a flame to them. The explosion which ensues is harmless. Plainly, to study the combination of hydrogen and oxygen by exploding a mix- ture of the gases will require special precautions. It can be carried out by the aid of the eudiometer (from evdia, good air, and /xerpov, a measure, an instrument for deter- mining the purity of air). The eudiometer is simply a tube graduated in millimetres and having two small plati- 74 COLLEGE CHEMISTRY. num. wires passed through it at the closed end, nearly meeting inside and ending in loops outside, as shown in Fig. 23. The eudiometer is filled with mercury, inverbed in a mercury trough, and held in an upright position by means of proper clamps. For the purpose of the experi- ment a quantity of hydrogen is brought into the tube, and ^lilULUUlWUliJlllUUlUIllltUUIUlUlUUUllllUllUlM Fig. £3. its volume accurately measured. About half this volume of oxygen is then introduced and the volume again accurately determined, and after the mixture has been allowed to stand for a few minutes a spark is passed between the wires in the eudiometer by connecting the loops with the poles of a small induction coil. Under these circumstances the explosion takes place noiselessly and with little or no danger. If the interior of the tube was dry before the explosion, it will be seen to be moist afterwards, and a marked decrease in the volume of the gases is also observed. That water is the product of the action has been proved beyond any possibility of doubt, over and over again. As the liquid water which is formed occupies an inappreciable volume as compared with the volume of the gases which combine, the decrease in volume represents the total volume of the hydrogen and oxygen which have combined. Now, if the experiment is per- formed with the two gases in different proportions, it will be found that only when they are mixed in the proportion of 2 volumes of hydrogen to 1 volume of oxygen do they completely disappear in the explosion. If there is a larger proportion of hydrogen present, the excess is left over; and the same is true of the oxygen. It will thus be seen that when hydrogen and oxygen combine to form water, they do so in the proportion of 2 volumes of hydrogen to 1 volume of oxygen, or more accurately 2.0027 to 1. Calculation of the Results Obtained in Exploding Mix- tures of Hydrogen and Oxygen. — Having determined that HYDROGEN AND OXYGEN— WATER-VAPOR. 75 whenever hydrogen and oxygen combine, they do so in the proportion 1 volume oxygen to 2 volumes hydrogen, and that when they combine the volume of liquid water formed measures so little as to amount to nothing in the measure- ments, we know that whenever a mixture of hydrogen and oxygen is exploded, no matter in what proportions they may he present, the volume of gas which disappears as such consisted of 2 volumes of hydrogen and 1 volume of oxygeD, or, in other words, one-third of the volume which disappears was oxygen and two-thirds hydrogen. Take this example: A quantity of hydrogen corresponding to 60 cc. under standard conditions is introduced into a eudiometer; 40 cc. of oxygen are added. What contrac- tion will there be on exploding the mixture ? Plainly the 60 cc. of hydrogen will combine with 30 cc. of oxygen. The 90 cc. of gas will disappear, and 10 cc. of oxygen will remain uncombined. From a total volume of 100 cc, therefore, we get a contraction to 10 cc. One-third of the contraction represents the oxygen and two-thirds the hydrogen. Determination of the Volume of Water-vapor formed by Union of Definite Volumes of Hydrogen and Oxygen. — The experiments which have been described enable us to draw the conclusion that hydrogen and oxygen combine in certain proportions by volume and by weight, and that a definite weight of water is formed; further, that the volume of liquid water formed when the two gases combine is inappreciable as compared with that of the gases. The question remains to be answered, What relation exists between the volumes of the combining gases and that of the water in the form of vapor ? This can be determined by causing the gases to combine in a eudiometer at a tem- perature sufficiently high to keep the water in the form of vapor. The simplest arrangement for accomplishing this is that shown in Fig. 24. A long eudiometer, the upper half of which is divided into three equal divisions marked on the outside, is filled with mercury and inverted in a bath of mercury. It is then surrounded by a large tube 7 6 COLLEGE CHEMISTRY. or jacket arranged as shown in Fig. 24. This is connected through the cork at the lower end with a vessel from which a current of steam can be obtained. The steam is passed through the jacket until the temperature of the mercury has reached that of the steam. A mix- ture of hydrogen and oxygen in the proportions in which they combine, viz., 2 of hydrogen to 1 of oxygen, is then introduced into the eudiometer so that it is filled to the third mark. This must be somewhat above the level of the mercury in the bath, so that the gases in the eudiometer shall be under diminished pressure. On now passing the spark the gases unite and the water which is formed re- mains in the form of vapor, as the temperature inside the eudiometer is nearly that of boiling water, and the vapor is under diminished pressure. It is found that the volume of the water-vapor is less than that of the gases introduced into the eudiometer. In order to secure the same pressure as that under which the gases were measured, the eudiometer must be lowered until the height of the mercury column in it is the same as it was before the explosion. On now measuring the volume of water-vapor, it will be found to be two-thirds that occu- pied by the uncombined gases. Therefore 2 volumes of hydrogen combine with 1 volume of oxygen and form 2 volumes of water-vapor. lb is an interesting fact that these simple relations exist between the volumes of the combining gases and the volume of the product. We shall see that similar relations hold good in the case of other gases; and the following general statement is based upon a great deal of careful study : When two or more gaseous substances combine to form a Fig. 24. HYDROGEN AND OXYGEN— HEAT. 77 gaseous compound, the volumes of the individual constit- uents as well as their sum bear a simple relation to the volume of the compound. This is known as the law of combination by volume. As will be seen further on, it has a most important bear- ing upon some of the fundamental ideas held in regard to the constitution of matter. Heat Evolved in the Union of Hydrogen and Oxygen. — To get as complete a knowledge as possible of the reac- tion which takes place between hydrogen and oxygen we have still to determine the amount of heat evolved. The heat evolved in burning a gram of hydrogen can be deter- mined, and from this we can calculate the heat of formation of water which, according to what was said on page 40, is the amount of heat evolved by the combination of 2 grams of hydrogen with 15.88 grams of oxygen. We are to determine the value of x in the equation . \\^ [H 2 , 0] = x cal., C\\e^ which expresses in thermochemical language the fact that when 2 grams of hydrogen combine with 15.88 grams of oxygen x calories of heat are evolved. The determination is made by burning a known weight of hydrogen in a vessel surrounded by water and arranged in such a way that all the heat is absorbed by the water. Experiment has shown that when 1 gram of hydrogen is burned 34,180 calories are evolved. Or, [H 2 , 0] = 68,360 cal. "So other substance gives as much heat as this in propor- tion to the weight used. Applications of the Heat formed by the Combination of Hydrogen and Oxygen. — To burn hydrogen in the air is, as we have seen, a simple matter, but to burn it in oxygen requires a special apparatus to prevent the mixing of the gases before they reach the end of the tube. where the com- bustion takes place. The oxyhydrogen bloivpipe answers 78 COLLEGE CHEMISTRY. this purpose. This may be constructed in several ways, but the simplest is that represented in Fig. 25. It con- sists of a tube through which a smaller tube passes. The Fio. 25, hydrogen is admitted through a and the oxygen through i. It will be seen that they come together only at the end of the tube. The hydrogen is first passed through and lighted; then the oxygen is passed through slowly, the pressure being increased until the flame appears thin and straight. It gives very little light but is intensely hot. Iron wire, steel, capper, zinc, and other metals burn in the flame with ease. Platinum vessels are made by melting the platinum by means of the oxyhydrogen flame. Oxyhydrogen Light. — When the oxyhydrogen flame is allowed to play upon some substance which it cannot melt or burn, the substance becomes heated so high that it gives off an intense light. The substance commonly used for this purpose is lime. Hence the light is often called the lime light. It is also known as the Drummond light. The hydrogen is first allowed to pass through the stop- cock, and lighted, when the oxygen is admitted. The flame plays against the piece of lime, and from this the light is given off when it has acquired a high temperature. Coal-gas may be and is generally used instead of hydrogen. Cylinders of compressed coal-gas and of oxygen can be bought, and the gases so prepared are used for the purpose of projecting pictures upon screens in illustrating lectures and for other similar purposes. Summary. — In our study of the action of hydrogen on oxygen, we have learned: (1) the relations between the weights of the two gases which act upon each other; (2) the relations between the volumes of the combining gases; COMPOSITION OF WATER. 79 (3) the relations between the volumes of the combining gases and that of the water- vapor formed; and (4) the amount of heat evolved when a given weight of hydrogen combines with oxygen. It remains for us to study more carefully the product formed. This is water. EXPERIMENTS. Composition ob "Water. * Experiment 47. — Arrange the apparatus shown in Fig. 22 ■with a straight tube instead of the bent tube E, and connect this with a small bent tube containing calcium chloride, as shown in Fig. 26. "Weigh tube E empty, and after the cupric oxide has been put into it. This gives the weight of the cupric oxide. Weigh the tube F before the experiment. Now proceed as in Experiment 45. In this case all the water formed by the action of the hydrogen on the cupric oxide will be absorbed by the cal- cium chloride in tube F. This tube will therefore gain in weight, and as oxygen is removed from the cupric oxide, tube E will Fig. 26. lose in weight. After the reduction is complete weigh tube E and tube F again. Let * = weight of tube E + cupric oxide before the experi- ment ; y = weight of tube E + copper after the experiment. Then x — y = weight of oxygen removed from the cupric oxide. Let a = weight of tube F before the experiment, and b = " " " after " " Then b — a = weight of water formed. If the experiment is properly performed, it will be found that the ratio =■ is very nearlv -t;- Or the result may be stated o — a ' 9 thus : In nine parts of water there are eight parts of oxygen. 8o COLLEGE CHEMISTRY. Experiment 48. — The tubes in the apparatus used in Experi- ment 3, or some other similar apparatus, should be graduated. Let the gases formed by the action of the electric current, as in Experiment 3, rise in the tubes, and observe the volumes. It will be seen that when one tube is just full of gas, the other, if it is of the same size, will be only half full. On examining the gases the larger volume will be found to be hydrogen, and the smaller volume oxygen. What are the relative weights of equal volumes of hydrogen and oxygen ? In what proportion by weight are the two gases obtained from water in this experiment ? How does this result agree with that obtained in the preceding experi- ment ? Does this experiment prove that water consists only of y hydrogen and oxygen ? \ Experiment 49.— Pass hydrogen from a generating-flask or a gasometer through a tube containing some substance that will absorb moisture, as all gases made in the ordinary way and collected over water are charged with moisture. The calcium chloride should be in granulated form, not powdered. After passing the hydrogen through the calcium chloride, pass it through a tube ending in a narrow opening, and set fire to it. If now a dry vessel is held over the flame, drops of water will con- dense on its surface and run down. A convenient arrangement of the apparatus is shown in Fig. 27. Fio. 27. A is the calcium chloride tube. Before lighting the jet hold a glass plate in the escaping gas, and see whether water is deposited on it. Light the jet before putting it under the bell-jar, other- EUDIOMETRIC EXPERIMENTS. 8 1 wise if hydrogen is allowed to escape into the vessel it will con-, tain a mixture of air and hydrogen, and this mixture, as we shall soon see, is "explosive. Experiment 50. — Mix hydrogen and oxygen in the proportion of about 2 volumes of hydrogen to 1 volume of oxygen in a gas- ometer. Fill soap-bubbles, made as directed in Experiment 30, with this mixture and allow them to rise in the air. As they rise, bring a lighted taper in contact with them, when a sharp explo- sion will occur. Great care must be taken to keep all flames away from the vicinity of the gasometer while the mixture is in it. This experiment is conveniently performed by hanging up, about six to eight feet above the experiment-table, a good-sized tin funnel-shaped vessel with the mouth downward. Now place a gas jet or a small flame of any kind at the mouth of the vessel. If the soap-bubbles are allowed to rise below this apparatus they will come in contact with the flame and explode at once.* What does this experiment show ? Does it give any informa- tion in regard to the composition of water ? Eudiometric Experiments. Experiment 51. — The general method of studying the combin- ation of hydrogen and oxygen by means of the eudiometer was described in the text (see p. 73). To what was there said it need only be added that, in exploding the mixture in the eudiometer, the latter should be held down firmly, by means of a clamp, against a thick piece of rubber cloth placed on the bottom of the mercury-trough. In making the measurements of the volume of the gases and the height of the mercury column, care must be taken to have the eudiometer in a perpendicular position. This can be secured by means of plumb-lines suspended from the ceil- ing and reaching nearly to the table, by which the position of the eudiometer can be adjusted. OXYHYDROGEN - BLOWPIPE. Experiment 52. — Hold in the flame of the oxyhydrogen blow- pipe successively a piece of iron wire, a piece of a steel watch- spring, a piece of copper wire, a piece of zinc, a piece of platinum wire. * The same apparatus may be used in experimenting with soap, bubbles filled with hydrogen. 88 COLLEGE CHEMISTRY. Experiment k 53. — Cut a piece of lime of convenient size and shape, say an inch long by three-quarters of an inch wide, and the same thickness. Fix it in position so that the flame of the oxy- hydrogen blowpipe will play upon it. The light is very bright, but by no means as intense as the electric light. CHAPTBK V. WATER. Historical. — Water was long thought to be an elementary substance until, towards the end of the last century, the discovery of hydrogen and oxygen, and of the nature of combustion, led to the discovery of its composition. Occurrence. — Besides the form in which water occurs in such enormous quantities in the earth, it also occurs in forms and conditions which prevent its immediate recogni- tion. Thus all living things contain a large proportion of water, which can be driven off by heat. If a piece of wood or of meat is heated, liquids pass off, and by purification these can be shown to consist mainly of water. The pro- portion of water in animal and vegetable substances is very great. If the body of a man weighing 150 pounds were to be put in an oven and thoroughly dried, there would be only about 40 pounds of solid matter left, most of the rest being water. Water also occurs in another form in which it does not directly reveal its presence. This is as water of crystal- lization. Many natural and manufactured chemical com- pounds give off water when heated. If, for example, the zinc sulphate formed in the preparation of hydrogen from zinc and sulphuric acid is dried by exposure to the air, or by pressing between layers of filter-paper, it will be found that, when heated in a dry tube, it gives off water, and at the same time changes its appearance. The same is true of gypsum which is found in nature, and of copper sul- phate or blue vitriol. In this last case the lose of water is 83 84 COLLEGE CHEMISTRY. accompanied by a loss of color. After all the water is driven off, the powder left behind is white. Many compounds when deposited from solutions in water in the form of crystals combine with definite quantities of water. This water is not present as such, but is held in chemical combination. Hence the substance does not appear moist, though it may contain more than half its weight of water. |TTiiH water of crys tallization i s. in some way which is not understood, essential to the form of the crystalsj If it is driven off, the crystals generally crumble to pieces. Some compounds combine with different quan- tities of water under different circumstances, the form of the crystals varying with the quantity of water held in combination. Compounds differ greatly as regards the ease with which they give up water of crystallization. In general, it is given off when the compound containing it is heated to the temperature of boiling water. But some compounds give it up at the ordinary temperature. This is true of sodium sulphate or Glauber's salt, which contains a quan- tity of water of crystallization represented by the formula Na 2 S0 4 . 10H 2 O. If some of the crystals are allowed to lie exposed to the air they undergo a marked change in the course of an hour or two. They lose their lustre and gradually crumble to pieces. [[Substances which lose their water of crystallization by exposure to the air are said to be efflorescent Some compounds if deprived of their water of crystal- lization will take it up again when allowed to lie in an atmosphere containing moisture, as, for example, in the air. It is well shown by the compound calcium chloride, OaCl 2 . If exposed to the air it will be noticed that it soon has a moist appearance, and if it is allowed to lie long enough it will dissolve in the water which it absorbs from the air. Substances which absorb enough water from the air to cause them to dissolve are said to be deliquescent. Formation of Water and Proofs of its Composition If we had not already learned, in studying the action of FORMATION OF WATER. 85 hydrogen upon oxygen, that water is composed of these two gases, we should first subject it to analysis. For this purpose we should have to bring it under the influence of a number of reagents and study its conduct. If we should pass an electric current through it in the proper way, we should observe that a gas rises from each pole. By placicg each pole under the mouth of an inverted tube filled with water the gases are easily collected. When one of the tubes has become full of gas the other one will be only half full. An examination will show that the gas in the tube which is filled is hydrogen, whereas that in the one which is half filled is oxygen. By decomposition of water by means of the electric current, then, there are obtained two volumes of hydrogen for each volume of oxygen. We already know the relative weights of equal volumes of the two gases, so that we can easily calculate the relative weights of the gases obtained in the experiment. The ratio of the weights of equal volumes of hydrogen and oxygen is 1 : 15.89. Therefore, if we have 2 volumes of hydrogen combined with 1 volume of oxygen, the ratio between the weights is 2. 0027 : 15. 89 or 1 : 7. 9 5. Although we know from the experiment referred to that hydrogen and oxygen are obtained from water in certain proportions, it^does not follow that this is the composition of water, f For it may be that other elements besides hydrogen and oxygen are contained in it, and it may be also that all the hydrogen and oxygen are not set free by the action of the lectric current. We might determine whether either of ihese possibilities is true or not by decomposing a weighed quantity of water, and weighing the hydrogen and oxygen obtained from it. If we should find that the sum of the weights of hydrogen and oxygen is equal to the weight of the water decomposed, this fact would be evidence that only hydrogen and oxygen are contained in water, and that they are present in the proportion stated. The same thing can be satisfactorily proved by causing hydrogen and oxygen to combine, or by effecting the synthesis of water. How this can be done has already been pointed out. It 86 COLLEGE CHEMISTRY. was shown, in the first place, that by burning hydrogen in oxygen water is formed. This proves that water con- sists of hydrogen and oxygen, but it does not furnish any proof as to the relation between the weights of the gases that combine. It is a qualitative synthesis. Other methods were described, the object of which was to show in what proportion by weight and by volume hydrogen and oxygen combine to form water. These methods are ex- amples of quantitative syntheses. The results proved that to form water hydrogen and oxygen combine in the pro- portion of 1 volume of oxygen to 2 of hydrogen; and it therefore follows that the decomposition of water which is effected by the electric current is complete. "-Properties of Water. — Pure water is tasteless and in- odorous, and in small quantities colorless. Thick layers are, however, blue. This is seen by filling a long tube with carefully purified water, and examining it by trans- mitted light, when it appears blue. Some mountain lakes also have a marked blue color. When cooled, water con- tracts until it reaches the temperature of 4°. At this point it has its maximum density. If cooled below this it expands, and the specific gravity of ice is somewhat less than that of water. Hence ice floats on water. If this were not so, in cold climates the water in the streams would freeze solid. As it is, the lower layers of water are protected by the ice and the cold water just below it, which are poor conductors of heat. Water can be cooled down below its freezing temperature, or 0°, if it is kept perfectly quiet, protected from the air, or cooled in capillary tubes. Water thus cooled down will suddenly solidify when dis- turbed, and then its temperature rises to 0°. Water boils at 100° under 760 mm. pressure. Increased pressure raises the boiling-point, and decreased pressure lowers it. Chemical Properties of Water. — Water is a very stable chemical compound. An indication of the great energy of the act of combination of hydrogen and oxygen is given in the great evolution of heat. In order to decompose water by heat, as much heat must be added as is evolved CHEMICAL PROPERTIES OF WATER. 87 when it is formed. At high temperatures it is decomposed into its elements. The decomposition begins at 1000° and is half complete at 2500°. On cooling the mixed gases they recombine. The amount of decomposition is constant for a given temperature. This kind of gradual decom- position of a compound by heat, followed by a recombina- tion on reversing the conditions, is called dissociation. It is a common phenomenon in chemistry, and farther on we shall have occasion to study it more fully. Water is decomposed by an electric current, when something is added to it to make it a conductor of electricity, and partly by contact with sodium and potassium at ordinary tem- peratures, and by iron and carbon at higher temperatures. It combines directly with a large number of substances in the form of water of crystallization, and with others to form definite chemical compounds called hydrates or hydroxides. Thus the oxides of potassium, K 2 0, and of sodium, Ka 2 0, combine with water with evolution of much heat to form the compounds potassium hydroxide, KOH, and sodium hydroxide, NaOH, which, it will be remem- bered, are also formed by the action of the elements potas- sium and sodium on water with liberation of hydrogen. The reactions between the oxides and water are represented by the equations K 2 + H 2 = 2KOH; Na 2 + H 2 = 2NaOH. Similarly, lime or calcium oxide, CaO, acts upon water with evolution of heat, as is observed in the process of slaking. The change is like that which takes place with potassium and sodium oxides, and is represented thus: CaO + H 2 = Oa0 2 H 2 [or Ca(OH)J. The product represented by the symbol Oa(OH) 2 is known as calcium hydroxide or slaked lime. In the same way barium oxide, BaO, forms barium hydroxide: BaO + H 2 = Ba0 2 H 2 [or Ba(OH)J. 88 COLLEGE CHEMISTRY. We shall meet with many examples of this kind of action in our study of chemical reactions, and we shall see that the hydroxides form two of the most important classes of compounds, known as acids and bases. The hydroxides of potassium, sodium, calcium, and barium are, for exam- ple, bases; while certain hydroxides containing sulphur, nitrogen, and carbon are acids, — such as sulphuric acid, S0 2 (OH) 2 , nitric acid, N0 2 (OH), and carbonic acid, CO(OH) 2 . It is not believed that water as such is con- tained in these hydroxides. Nevertheless, when heated many of them give off water. Thus, when heated to a red heat, calcium hydroxide is decomposed into the oxide and water according to the equation Ca(OH) 2 = CaO + H 2 0. - Many substances that contain hydrogen and oxygen act in the same way. This is due to the great stability of water even at elevated temperatures. As the temperature becomes higher and higher the constituents of the com- pound are held together less and less firmly. When a point is reached at which the attraction of the hydrogen for the oxygen is greater than that required to hold the con- stituents together, a rearrangement takes place, and com- pounds which are stable at the higher temperature are formed. Water as a Solvent. — With a great many substances water forms unstable compounds, the nature of which cannot at present be explained. These unstable com- pounds are called solutions. It is known that many solids,' liquids, and gases when brought into water disappear and form colorless liquids which look like water. Some give colored liquids of the same color as the substance dissolved, and others give liquids which have a color quite different from that of the substance dissolved. On the other hand, there are many substances which do not form such com- pounds with water or which, as we say, are insoluble in water. In a solution the particles of the substance dis- WATER AS A SOLVENT. 89 solved are in some way held in combination by the particles of the liquid. If a very small quantity of substance is dis- solved in a large quantity of water and the solution thoroughly stirred, the dissolved substance is uniformly distributed throughout the liquid, as can be shown by refined chemical methods. That the dissolved substance is everywhere present in the solution can be shown further by the aid of certain dye-stuffs, as, for example, magenta. A drop of a concentrated solution of the substance brought into many gallons of water imparts a distinct color to all parts of the liquid. An experiment of this kind gives some idea of the extent to which the division of matter can be carried. For it is evident that in each drop of the dilute solution there must be contained some of the color- ing matter, though the quantity must be what we should ordinarily speak of as infinitesimal. While there seems to be no limit to the extent to which a solution can be diluted and still retain the dissolved substance uniformly dis- tributed through its mass, there is a limit to the amount of every substance that can be brought into solution, and this varies with the temperature, and, in the case of gases, with the pressure. Some substances are easily soluble, others are difficultly soluble. When the solutions are boiled the water simply passes off and leaves the dissolved sub- stance behind, if it is a non-volatile solid. \lf, noweveTT" the substance in solution is a liquid, a partial separation will take place, the extent of the separation depending largely upon the difference between the boiling-points of the water and the other liquid. A complete separation of two liquids by boiling is difficult and in most cases im- possible. If, finally, the substance in solution is a gas, it generally passes off when the solution is heated, though in some cases water is given off leaving the gas in solution, which of course then becomes more concentrated. When a certain concentration is reached a solution of the gas passes over. It is probable that in these cases the gas is in a condition of chemical combination with the water. Solutions in general seem to differ from true chemical $o COLLEGE CHEMISTRY. compounds in some important particulars, and also from mere mechanical mixtures. Definiteness of composition appears to be characteristic of chemical compounds, or, at least, it is characteristic of a large number of compounds which we call chemical compounds. But solutions have no definite composition. We can dissolve any quantity of a substance from the minutest particle to a certain fixed quantity, and the solutions formed are uniform and appear to be. just as truly solutions as that which contains the largest quantity which can be held in combination. On the other hand, in a mere mechanical mixture the constit- uents may be present in all proportions, while this is not true of solutions. The subject of solution is under inves- tigation, and it has been made highly probable that some substances, when dissolved, are partly or wholly broken down into smaller parts charged with positive and negative electricity respectively. These smaller parts are called ions. "Solution as an Aid to Chemical Action. — When it is /desired to secure the chemical action of one solid substance upon another, it is generally necessary to bring them together in solution. One reason why they do not act readily when mixed in the solid condition is to be found in the fact that, under these circumstances, their particles V: remain separated by sensible distances, no matter how finely the mixture may be powdered. If, however, the substances are dissolved, and the solutions poured together, the particles of the liquid move so freely among one another that they come in intimate contact, thus facilitating chemical action. Many substances which do not act upon one another at all when brought together in dry condition act readily when brought together in solution. It is believed that this is due principally to the splitting of the compounds into ions by the action of the water. These ions are free to act upon other ions which may be brought into the same solution. This idea will be developed farther on in other connections. Although it is highly probable, then, that when a reaction takes place in a water NATURAL WATERS. Qt solution the water itself plays a very important part, the reaction is generally represented by an equation in which the water does not appear. Of course, such an equation is imperfect, but it answers certain purposes quite satisfac- torily, and may be used without danger of confusion. Thus, when hydrochloric acid acts upon zinc, hydrogen is liberated and zinc chloride is formed. What we call hydrochloric acid in the laboratory is the liquid which is formed by the absorption of hydrochloric acid gas, HC1, by water. When this liquid is used, however, the chem- ical act which makes itself known to us is that which gives hydrogen gas and zinc chloride in solution. Apparently this reaction is independent of the water, and it may be represented thus : Zn + 2HC1 = ZnCl 2 + 2H. Sometimes such reactions are written as follows in order to express the fact that water is present, though, as will be observed, no attempt is made to tell what part the water plays : Zn + 2HC1 + Aq = ZnCl 2 -fc- 2H + Aq; or Zn + 2HC1 + H 2 = ZnCl 2 + 2H + H 2 0. There is no objection to this, certainly, but it is question- able whether, considering the purposes for which chemical equations are generally used, this increases their value. Natural Waters. — All water found in nature is more or less impure or, in other words, contains something in solution. In the first place, waters which are exposed to the air dissolve some of the gases of which the air is com- posed, as oxygen, nitrogen, and carbon dioxide. Again, natural waters necessarily are in contact with the earth; they always dissolve some of the earthy substances; and, finally, many waters come in contact with animal and vegetable substances and dissolve something from these. The water which is carried up as vapor from the surfaces of natural bodies of water is approximately pure. When this is precipitated as rain it dissolves certain substances 92 COLLEGE CHEMISTRY. from the air, and the first rain that falls during a storm is always more or less contaminated. In a short time, how- ever, the air becomes washed and the rain which falls thereafter is approximately pure water. If it remains in contact with insoluble rocks, as, for example, quartzite or sandstone, it remains pure, and mountain-streams which flow over sandstone beds are, in general, the purest. Water which flows over limestone dissolves some of this and becomes " hard." A similar change is brought about in water by contact with gypsum and magnesium sulphate. The condition of hardness will be taken up more fully under calcium and magnesium compounds. The many varieties of mineral springs have their origin in the pres- ence in the earth of certain substances which are soluble in water. Among those most frequently met with in solution in natural waters are carbonic acid, sodium car- bonate, sodium sulphate or Glauber's salt, sodium chloride or common salt,, magnesium sulphate, iron carbonate, and sulphuretted hydrogen. Effervescent waters are those which contain a large quantity of carbonic acid in solution and give off carbon dioxide gas when exposed to the air. Chalybeate waters are those which contain some compound of iron in solution; sulphur waters contain the gas, sul- phuretted hydrogen. Common salt occurs in large quantities in different parts of the earth. As it is easily soluble in water, many streams contain it; and as most streams find their way to the ocean, we see one reason why the water of the ocean should be salt. /"As streams approach the habitation of man they are /subjected to a serious cause of contamination. The ; drainage from the neighborhood of human dwellings is I very apt to find its way into a near stream. The sub- \ stances thus carried into the stream undergo decomposition and give rise to the formation of larger or smaller quanti- [ ties of new products, some of which have the power to i produce disturbances when taken into the system, and others to produce disease. This condition of things is most strikingly illustrated by the case of a large town PURIFICATION OF WATER. 93 situated on the banks of a river. It frequently happens that the water of the river is used for drinking purposes, and it also frequently happens that the water is contami- nated by drainage. Eiver water when once contaminated by drainage tends to become pure again by contact with the air, the change consisting largely in the slow oxidation of the substances which are of animal or vegetable origin, and their conversion into harmless products. If water is to be used for drinking purposes, however, it is not well to rely too much upon this process of purification. So much has of late years been said about drinking-water that excessive alarm has been created, and water is no doubt frequently held responsible for sickness with which it has nothing to do. In some places the war against the water- supply has been carried so far that those who can afford it drink only artificially purified and distilled water. It is undoubtedly well to be cautious, but it is possible to be too cautious. What Constitutes a Bad Drinking-water A good drinking-water should be free from odor and taste and should not contain anything that can act injuriously upon the system. It is, however, difficult to decide by chemical means whether the water contains anything injurious or not, as there may be a very minute quantity of an ex- tremely injurious substance, for example a disease germ, present, and chemical analysis would be powerless to detect it. On the other hand, water which is very considerably contaminated by sewage may be harmless, and yet the latter -might be pronounced "bad" and the former "good." The rule generally adopted by chemists in deal- ing with water is to pronounce any water dangerous that is contaminated by sewage. Such contamination can generally be detected by analysis or by analysis and inspec- tion of the sources. Purification of Water. — Impure water may become purer by natural methods as has been stated, and it- may be rendered fit for drinking purposes by filtering through such substances as charcoal, sand, spongy iron, etc. A 94 COLLEGE CHEMISTRY. filter, no matter of what it may be made, will not, how- ever, remain efficient for any length of time, as the sub- stances contained in the impure water are retained by it and, after a time, it becomes a source of pollution instead of a purifier. For refined work in chemistry pure water is prepared by distilling natural waters. The process of distillation consists in boiling the water and then passing the steam through a tube or system of tubes surrounded by cold water. Thus the steam is condensed, and the dis- tilled water is approximately pure. Of course, it is neces- sary that the tubes in which the condensation takes place should be of such material that water does not act upon it to any extent. The materials used are glass, block tin, and platinum. Chemically pure water is a very rare sub- stance even in the best chemical laboratories. The slight impurities present in ordinary distilled water are not, however, of special importance under ordinary circum- stances. EXPERIMENTS. Organic Substances contain Wateb. Experiment 54. — In dry test-tubes heat gently various organic substances as a piece of wood, fresh meat, fruits, vegetables, etc. "Water of Crystallization. Experiment 55. — Take some of the crystals of zinc sulphate obtained in Experiment 34. Spread them out on a layer of filter- paper, and finally press two or three of them between folds of the paper. Examine them carefully. They appear to be quite dry, and in the ordinary sense they are dry. Put them in a dry tube and heat them gently, when it will be observed that water con- denses in the upper part of the tube, while the crystals lose their lustre, becoming white and opaque, and at last crumbling to powder. Experiment 56. — Perform a similar experiment with some gypsum, which is the natural substance from which "plaster of Paris " is made. PURIFICATION OF IVATER. 95 Experiment 57. — Heat a few small crystals of copper sulphate, or blue vitriol. In this case the loss of water is accompanied by a loss of color. After all the water is driven off, the powder left behind is white. On dissolving it in water, however, the solution will be seen to be blue ; and if the solution is evaporated until the substance is deposited, it will appear in the form of blue crystals. Efflorescent Salts. Experiment 58.— Select a few crystals of sodium sulphate which have not lost their lustre. Put them on a watch-glass, and let them lie exposed to the air for an hour or two. They soon lose their lustre, and undergo the changes noticed in heat- ing zinc sulphate. Deliquescent Salts. Experiment 59.— Expose a few pieces of calcium chloride to the air. Its surface will soon give evidence of the presence of moisture, and after a time the substance will dissolve in the water which is absorbed. Purification of Water by Distillation. Experiment 60. — In an apparatus like that shown in Fig. 28 Fig 28. distil a dilute solution of copper sulphate or some other colored substance. A slow current of cold water must be kept running 96 COLLEGE CHEMISTRY. through the condenser by connecting the lower rubber tube with a water-cock. When the water is boiled in the large flask, the steam passes into the inner tube of the condenser. As this is surrounded by cold water, the steam condenses and the distilled water collects in the receiver. CHAPTER VI. CONSTITUTION OF MATTER. — ATOMIC THEORY.— ATOMS AND MOLECULES. — CONSTITUTION. — VA- LENCE. The Atomic Theory as proposed by l5alton. — We have seen how Dalton, at the beginning of this century, dis- covered the law of multiple proportions. This law, as well as that of definite proportions, called for explanation. The questions to be answered are: (1) Why do the elements combine in definite proportions ? (2) Why, when ele- ments combine with each other in more than one way, do the relative quantities which enter into combination in the different cases bear simple relations to one another ? and (3) What is the significance of the figures representing the combining weights ? Dalton saw that the facts referred to could be explained by the atomic theory, while, on the theory that matter is indefinitely divisible, they appear to be inexplicable. It is only necessary to assume that each element is made up of particles which are not divisible in chemical processes, and that these particles, or atoms, have definite weights. The atoms of any one element must be supposed to have the same weight, while the atoms of different elements have different weights. Now, when chemical combination takes place, Dalton supposed that the action is between the atoms. The simplest case is that in which combination takes place in such way that each atom of one element combines with one atom of another; but, besides this kind of combination, we may have that in which one atom of one element combines with two atoms of another, or two of one may combine with three of another, etc. Suppose two elements A and B, the 97 98 COLLEGE CHEMISTRY. weights of whose atoms are to each other as 1 : 10, are brought together, and they combine in the simplest way, i.e., one atom of one with one atom of the other, then it is plain that in the compound AB the elements will be contained in the proportion of 1 part of A to 10 parts of B, whether a small or a large quantity of the compound is formed, and no matter in what proportions the elements are brought together. If they should be brought together in the proportion of their atomic weights (1 : 10), then no part of either element will be left uncombined after the act of cdmbination has taken place. If, however, a larger proportion of either element is taken than that stated, then the quantity of the one which is in excess of this pro- portion will be left uncombined. This is in accordance with what we know takes place, and it is a conclusion drawn from the theory. No matter how many atoms of A we may take, the same number of atoms of B will be required to combine with all of them. But each atom of B weighs 10 times as much as each atom of A, therefore the total mass of B which enters into combination must be 10 times that of A with which it combines. It may be, however, that these same elements can form other compounds with each other. If A and B represent the atoms of the elements and we assume these atoms to be chemically indivisible, then the other compounds must be represented by such symbols as AB 2 , A 2 B, AB S , A 3 B, A 2 B S , etc., which represent compounds in which 1 atom of A is combined with 2 atoms of B; 2 of A with 1 of B; lot A with 3 of B; 3 of A with 1 of B; 2 of A with 3 of B; etc.; or they also represent compounds in which 1 part by weight of A is combined with 20 parts by weight of B; 2 of A with 10 of B; 1 of A with 30 of B; 3 of A with 10 of B; 2 of A with 30 of B; etc. It is therefore clear that, if the atomic theory as put forward by Dalton is true, the elements must combine according to the laws of definite and multiple proportions; and it appears that the figures which represent the combining weights of the elements must either bear to one another the same relation USE AND VALUE OF A THEORY. 99 as the weights of the atoms, or the atomic weights or the relative weights of the atoms must be closely related to the combining weights, as will be pointed out more clearly presently. Use and Value of a Theory.— The relation of a theory to facts is very simple, but is frequently misunderstood. The relation may be conveniently illustrated by the case under consideration. By a careful investigation of a number of chemical compounds it was shown that in each of them the same elements always occur in the same proportion. This led to the belief that this is true of every chemical compound, and after further investigation which, as far as it went, showed the surmise to be correct, the law of definite proportions was proposed. This law is" simply a statement of what has been found true in all cases examined. It involves no speculation. It is a statement of fact. It may be said that the statement or law must be open to some doubt for the reason that all possible cases have not been examined, and it may not hold true for some of these unexamined cases. The reply to this is that it has been found true in a very large number of cases and in all cases which have been investigated. It is true for the present state of our knowledge, and that is all we can demand of any law. Again, further investigation led to the discovery of the law of multiple proportions, which is also a statement of what has been found true in all cases investigated. It, like the law of definite proportions and in the same sense, is a statement of fact. But having gone thus far, we now ask, what is the explanation of these laws ? We know the facts — what is the explanation ? By experiment we cannot go beyond these facts, but it is possible to imagine a cause and then proceed to see whether the imagined cause is sufficient to account for the facts. This is what Dalton did. He imagined that matter is made up of atoms of definite weights, and that chemical combination takes place in simple ways between these atoms. This imagined cause is the atomic theory. It is not a statement of anything found by investigation. loo COLLEGE CHEMISTRY. It is not a statement of an established fact. It may or may not be literally true, but at all events it is the best guess that has ever been made as to the cause of the fundamental laws of chemical action, and it furnishes a very convenient means of interpreting the facts of chemis- try. Since the atomic theory was first proposed it has been accepted by nearly all chemists. It has been of great value ■ n suggesting methods of work, and has contributed largely to the advance of chemistry. Any theory which is in accordance with the facts and leads to the discovery of new facts is of value, whether it should eventually prove to be true or false. At the same time a false theory may do much harm, as it may lead men to misinterpret the facts which they observe, and thus retard progress. Atomic Weights and Combining Weights. — If the atomic theory is true, the atoms of each element must have definite weights, and the determination of these atomic weights must evidently be of great importance. By analy- sis of compounds we can only determine the proportions by weight in which the elements combine with one another. Can we in this way determine the atomic weights ? In the first place, it is clear that it is out of the question to think of determining the absolute weights of the atoms, and all that we can possibly do is to determine their relative weights. As of all the elements hydrogen enters into combination in smallest relative quantity, its atomic weight is taken as the unit of the system, and the problem before us is to determine how many times heavier the atoms of the other elements are than that of hydrogen. If every element combined with hydrogen in only one proportion the problem would be a comparatively simple one. Thus the three elements chlorine, bromine, and iodine combine with hydrogen, forming only one compound each. On analysis these are found to contain respectively 1 part of hydrogen to 35.18 parts of chlorine; 1 " " " 79.36 " " bromine; and 1 " " " 125.90 " " iodine. ATOMIC WEIGHTS AND COMBINING WEIGHTS. 101 There is no reason for believing that in these compounds the elements are combined in any but the simplest way, i.e., that each atom of hydrogen is combined with one atom of chlorine to form hydrochloric acid, etc. If this is true, then the atom of chlorine must weigh 35. 18 times, that of bromine 79.36 times, and that of iodine 125.90 times as much as that of hydrogen, or, in other words, the atomic weights of chlorine, bromine, and iodine are respectively 35.18, 79.36, and 125.90. It will, however, be observed that there is no evidence as to whether the elements in these compounds are combined in the simplest way or not. It is possible, as far as we know, that one atom of hydrogen may combine with two or three of chlo- rine, or that one of chlorine may combine with two or three of hydrogen. As there is, however, no evidence upon this point the simplest assumption is made. If we take the case of oxygen the problem is more com- plex. In water the elements are combined in the propor- tion of 7.95 parts by weight of oxygen to 1 of hydrogen, and from this we should naturally conclude that the atomic weight of oxygen is 7.95; but further study shows that this conclusion is not justified. Hydrogen and oxygen form a second compound known as hydrogen dioxide or hydrogen peroxide in which there are 15.90 parts by weight of oxygen to 1 of hydrogen. This may be explained in the terms of the atomic theory by assuming that water is represented by the formula HO, and hydrogen dioxide by H0 2 . But it may be that the atomic weight of oxygen is 15.89, and then water must be represented by the formula H 2 0, and hydrogen dioxide by HO. Simple analysis of the compounds is not sufficient to enable us to decide between these possibilities. It is, therefore, evident that in order to determine the atomic weights something besides the determination of the composition of compounds is necessary. The figures representing the combining weights found in this way will, however, either be identi- cal with the atomic weights or will bear a simple numerical relation to them. 102 COLLEGE CHEMISTRY. Molecules. — Investigation of certain phenomena of light, of electricity, of liquid films and the conduct of gases has led physicists to the conclusion that matter is not continuous, but made up of small particles, which are called molecules. A gaseous molecule is defined as "that minute portion of a substance which moves about as a whole, so that its parts, if it has any, do not part company during the motion of agitation of the gas." It would be out of place here to present the physical facts upon which the molecular theory rests. Suffice it to say that it is the only theory which has been found adequate to account for the behavior of gases. Avogadro's Hypothesis.— The fact that gases conduct themselves in the same way under the influence of changes in temperature and pressure can only be explained by assuming that equal volumes of all gases and vapors contain the same number of ultimate particles or molecules at the same temperature and pressure. This is a deduction from the well-tested molecular theory of gases. It was, however, originally put forward by the Italian chemist Avogadro from a study of chemical as well as of physical facts, and a little later it was sug- gested as probable by the French physicist Ampere. It is therefore generally spoken of as Avogadro's hypothesis, and sometimes as Ampere's hypothesis. Absolute proof of its truth cannot be given, but it is in thorough accord- ance with a large number of well-known facts, and it is undoubtedly true, if the molecular theory of matter is true. It may therefore be considered as furnishing a solid foun- dation for further conclusions bearing upon the problem of the determination of the atomic weights. Distinction between Molecules and Atoms. — If we con- sider any chemical compound, as water or hydrochloric acid, it is evident that the smallest particle or the molecule of the compound must be made up of still smaller particles. Thus, the smallest particle of water must contain smaller particles of hydrogen and oxygen, and the smallest particle of hydrochloric acid must contain smaller particles of MOLECULAR WEIGHTS. 103 hydrogen and chlorine. These smallest particles of the molecules are the atoms. The molecules of the compounds are, according to this view, made up of the atoms of the elements. Similarly the elements themselves are, for good reasons which will be presented, believed to consist of molecules which are in turn made up of atoms of the same kind, though in a few cases the molecule of the element is identical with the atom. The difference between a com- pound and an element then is, in general, that the molecule of the compound consists of atoms of different kinds, while the molecule of an element consists of atoms of the same kind or, in a few cases, of one atom. Generally the atoms do not exist in the free or uncombined state, but, if they are set free by chemical action, they unite to form mole- cules. The following may serve as a definition of the con- ception of atoms at present held by chemists : Atoms are the indivisible constituents of molecules. They are the smallest particles of the elements that take part in chemical reactions, and are, for the greater part, incapable of existence in the free state, being generally found in com- bination with other atoms, either of the same hind or of different kinds. It cannot be too strongly emphasized that the views held in regard to the relations between molecules and atoms are based upon an enormous amount of painstaking study of facts, and in order fully to comprehend their value a study of most of these facts would be necessary. These views have gradually become firmly established as knowledge of the facts has grown more and more profound. Accepting them, we are now to see how they aid us in the problem with which we are dealing, viz., the determination of the atomic weights. Molecular Weights. — If equal volumes of gases contain the same number of molecules at the same temperature and pressure, it is only necessary to determine the weights of equal volumes of gases to learn the relative weights of their molecules. Thus, if we weigh a litre of each of three gases, and find that the weights are to one another as 1 to 104 COLLEGE CHEMISTRY. 2 to 3, then it follows that the relation between the weights of the molecules of these gases is expressed by these figures, or, in other words, the molecule of the second gas is twice as heavy, and that of the third gas is three times as heavy as that of the lightest. The determination of the relative weights of the molecules of substances which either are gaseous or can be converted into gases resolves itself simply into a determination of the weights of equal volumes. In representing the molecular weights we may use any figures that are most convenient, provided only that they bear to one another the relations determined by experiment. If, however, we call the atomic weight of oxygen 16, then our system of molecular weights must be based upon this standard. But 16 is not the molecular weight of oxygen. In other words, the atom and the molecule of oxygen are not identical. This conclusion is drawn from a large number of facts. It is the result of a half century's work by the leading chemists of the world. One way in which it can be made clear is this : The weight of a given volume of oxygen gas as compared with the weight of the same volume of water in the form of gas under the same conditions of temperature and pressure is as 32 to 18.02. These equal volumes of gases contain the same number of molecules according to the hypothesis of Avogadro. Bach molecule of water is made up of 1 atom of oxygen with the atomic weight 16 and 2 atoms of hydrogen with the atomic weight 1.01. If 16 is the atomic weight of oxygen, water cannot have a smaller molecular weight than 18.02. But, if the molecular weight of water is 18.02, that of oxygen must be 32, or it must be twice as great as the atomic weight. No matter what atomic weight we may adopt for oxygen this reasoning holds good. The " standard molecule is therefore that of oxygen or, better, the basis of the system is a hypothetical gas having -£% the weight of oxygen. This may be called the normal gas. Having adopted this standard, the molecular weight of a gas may be defined as its weight compared with that of the normal gas. Below are given the molecular weights DEDUCTION OF ATOMIC WEIGHTS. i°5 of a few elements and compounds which have been deter- mined by the method just described : Name. Hydrogen Nitrogen Water Hydrochloric acid Ammonia Marsh-gas , Carbon monoxide. Carbon dioxide. . . Molecular Molecular Weight. Formula. 2.02 H„ 28.08 N a 18.02 H 3 36.46 HC1 17.07 NH 8 16.04 CH« 28 CO 44 CO a Deduction of Atomic Weights from Molecular Weights. — The determination of molecular weights does not neces- sarily carry with it the determination of the atomic weights. It is plain from what has already been said that a knowledge of the molecular weight of an element does not convey a knowledge of its atomic weight. If, for example, we learn that the molecular weight of nitrogen is approximately 28, we have no means of judging from this what the atomic weight is. It is plainly necessary to know of how many atoms each molecule of nitrogen is made up, and to learn this is not a simple matter. It is easier to determine the atomic weight of an element through a study of its compounds. Suppose it is desired to determine the atomic weight of oxygen. We first determine the molecular weights of a number of com- pounds that contain oxygen, and then analyze these com- pounds. We then see what the smallest figure is that is required to express the weight of the oxygen that enters into the composition of the molecules, and this figure is selected as the atomic weight. The molecular weights and the composition of several oxygen compounds are given in the following table. The figures in the third column are of course determined by analysis, an example of the methods used having been given in the chapter on water. Stated in ordinary language, the figures in the case of carbon mon- io6 COLLEGE CHEMISTRY. Composition. 2.02 parti 3 hydrogen, 16 it oxygen. 12 " carbon, 16 a oxygen. 12 it carbon, 32 a oxygen. 14.04 ti nitrogen, 16 a oxygen. 28.08 a nitrogen, 16 1 1 oxygen. 32.06 " sulphur, 32 it oxygen. 32.06 a sulphur, 48 a oxygen. Compound. Mol. Wt. Approx. Water 18.02 Carbon monoxide 28 Carbon dioxide 44 Nitric oxide 30.04 Nitrous oxide 44.08 Sulphur dioxide 64.06 Sulphur trioxide 80.06 oxide mean that the molecule of this compound weighs 28 times as much as the normal gas, and the 28 parts by weight of matter are made up of 12 parts of carbon and 16 parts of oxygen. Considering now the composition of the com- pounds in the table, it will be seen that the smallest mass of oxygen that enters into the composition of any of the molecules weighs 16 in terms of the normal gas. We find twice this mass, as in carbon dioxide and sulphur dioxide; and three times, as in sulphur trioxide, but no smaller mass. Now, an examination of all compounds of oxygen that are known to us in the form of gas or vapor shows the same thing to hold true; that is to say, the smallest mass of oxygen that enters into the composition of mole- cules is 16 times as great as that of the normal gas. The conclusion is therefore drawn that 16 is the atomic weight of oxygen. The possibility that the atomic weight of oxygen is less than this figure is not excluded. It may be that in the simplest oxygen compounds now known there are two or more atoms of this element in the molecules. But in the total absence of evidence on this point all we can do is to accept the figure 16 as in perfect accordance with all our knowledge of oxygen compounds. In this way the atomic weights of all elements that form gaseous compounds or compounds that can be converted MOLECULAR FORMULAS, 107 into vapor have been determined; and the determinations made in this way are regarded as the most reliable. Exact Atomic Weights determined by the Aid of Analysis. — By determining molecular weights it is possible to decide approximately what figure represents the atomic weight of an element, but the methods employed in mak- ing determinations of molecular weights are liable to slight errors, and therefore the atomic weights obtained directly from the molecular weights deviate slightly from the true figures. In order to determine the atomic weights with the greatest possible accuracy, the most refined methods of chemical analysis are brought into play, and the figures in the table on page 19 have been determined in this way by a combination of a study of the specific gravity of gases and by the most careful analyses, together with some other methods which will be taken up later. Molecular Formulas. — The symbols of chemical com- pounds first used were intended to express simply the composition of the compounds, and this can be done as was explained in Chapter I, by adopting a system of com- bining weights of the elements. According to the theory explained in the last chapter, the smallest particle of every compound is a molecule, and each molecule is made up of atoms. It appears, therefore, desirable for the sake of uniformity that the symbols used to represent chemical compounds should represent molecules. Where the molec- ular weight of a compound, the atomic weights of the elements of which it is composed, and its composition are known, there is no difficulty in representing it by a molec- ular formula. Thus, the molecular weight of ammonia is found by experiment to be approximately 17, and the 17 parts are made up of 14 parts of nitrogen and 3 parts of hydrogen. The atomic weight of nitrogen is found by the method which has just been described to be very nearly 14. Therefore the molecule of ammonia weighing 17 parts is composed of 1 atom of nitrogen weighing 14 parts and 3 atoms of hydrogen weighing 3 parts. The compo- sition of the molecule is therefore represented by the 108 COLLEGE CHEMISTRY formula NH 3 . Similarly the composition of the molecule of water is represented by the formula H 2 0; that of hydro- chloric acid by HOI; that of marsh-gas by CH 4 ; etc., etc. Every formula now in use is intended to represent a mole- cule of the compound for which it stands. In regard to the molecular weights of compounds that are not gaseous nor convertible into vapor, Avogadro's method is plainly of no avail. Methods have, however, been devised which are applicable to a number of these (see Chapter XXIII). Constitution. — When hydrochloric acid is formed, we conceive that each atom of hydrogen combines with one atom of chlorine, and that the molecules of the resulting compound are made up each of an atom of hydrogen and an atom of chlorine. What the act of combination con- sists in we do not know. We simply know that something very remarkable takes place, and that as a consequence the hydrogen and chlorine cease to exist in their original forms. It is idle at present even to speculate in regard to the character of the change. The fact of union is ex- pressed by writing the symbols of the elements side by side without any sign between them, as HC1, or, sometimes, it is convenient to use a line to indicate chemical union, thus: H-Cl. According to the molecular theory the molecule of water consists of two atoms of hydrogen and one of oxygen, as represented by the formula H 2 0, and the question now suggests itself whether all three atoms are in combination with one another or whether each of the hydrogen atoms is- in combination with the oxygen atom, but not with each other, as represented by the formula H-O-H. So, too, in the case of ammonia, the molecular formula of which is NH 3 , the question suggests itself : Are the three atoms of hydrogen in combination with the atom of nitrogen, but not with one another, as represented in the formula N^-H? It is extremely difficult to answer such questions, but, nevertheless, certain facts are known which enable us to draw probable conclusions. Formulas VALEXCE. 109 which express the composition of molecules and at the same time express the relations or the connections which exist between the atoms are called constitutional formulas. These constitutional formulas are very frequently used at present, but sometimes without a sufficient basis of facts to justify them. Whenever they are used in this book, the reasons for them will be stated as fully as may appear necessary. Valence. — The formulas of the hydrogen compounds of chlorine, oxygen, nitrogen, and carbon, all determined by the same method, are C1H OH s NH S CH 4 . A consideration of these formulas and of many similar ones has led to the belief that the atoms of different ele- ments differ in their power of holding other atoms in combination. The simplest explanation of the composition of the compounds above represented is that the atoms of chlorine, oxygen, nitrogen, and carbon differ in their power of holding hydrogen atoms in combination. Hy- drogen and chlorine combine in only one way, 1 atom of chjprine combining with 1 of hydrogen; 1 of oxygen com- bines with 2 of hydrogen; 1 of nitrogen with 3 of hydrogen; and 1 of carbon with 4 of hydrogen. The limit of the combining power of the atom of chlorine is reached when it has combined with one atom of hydrogen. And as one chlorine atom can hold but one atom of hydrogen in combination, so one atom, of hydrogen can hold but one atom of chlorine. Either the hydrogen atom or the chlorine atom may be taken as an example of the simplest kind of atom. Any element like hydrogen or chlorine is called a univalent element ; an element like oxygen whose atom can hold two unit atoms in combina- tion is called a bivalent element j an element like nitrogen whose atom can hold three unit atoms in combination is called a trivalent element ; and an element like carbon whose atom can hold four unit atoms in combination is called a quadrivalent element. Most elements belong to no COLLEGE CHEMISTRY. one or the other of these four classes, though there are some which can hold five, six, and even seven unit atoms in combination. These are called quinquivalent, seziva- lent, and septivalent respectively. Valence is defined as that property of an element by virtue of which its atom can hold a definite number of other atoms in combination. In the formation of com- pounds the valence of the elements determines how many atoms of any element can enter into combination with any other. The atoms are sometimes spoken of as having bonds which are graphically represented by lines. Thus, a univalent element is said to have one bond, as repre- sented by II-, C1-, etc. ; a bivalent element is said to have two bonds, -0-, -S-, etc. ; a trivalent element three, -N-; i and a quadrivalent element four, -0-. Of course, this is merely a symbolical representation of the idea that each atom has a definite power of combining with others. It is further said that when the atoms unite these bonds become satisfied. Thus when one atom of hydrogen unites with one of chlorine, the bond of each is regarded as uniting with the bond of the other, and this is represented by the symbol H-Cl. So, too, when two atoms of hydrogen unite with one of oxygen, the compound is represented in this way: H-O-H or 0 Other groups are those of phosphorus, 31, vanadium, 51.2, and arsenic, 75: il±15 =6 3; lithium 7.03, sodium, 23.05, and potassium, 39.15: 7.03 + 39.15 2 = 23.09. Arrangement of the Elements Mendeleeff (1869) and Lothar Meyer (1870) have proposed several arrangements * A large table of the Natural System of the Elements should be hung up in a conspicuous place in the laboratory. 173 174 COLLEGE CHEMISTRY. 00 OS "ft-: «?to II 3 s i-i a CD CO ■ 5 2°. "illl 1 "iiU 1 > 1 O 3 U) i y 1 I « 1-T S«3 6o"oo" oo CJ II II II II 1 1 CD O ^■o ho PSft< 1 ta. *~*?~' o ■~ 1r o» 1 ''■*■"' l i-H IO CD S9 1 1 P w" q en II m - II " g II i A II ° fl II t-l O h s 1 i 1 1 C4 en « 1 | P I" s C3 B 03 i CO II o CO II ™ CO » II u O «- en " & & 11 o a in II i 1 00 II 1 en 1 II u ,_ 10 CO 1 CO p i 3 o II CO II a* II ii M - 11 -0" to OS o CM II to II 1 00 00 II 8 ii 3 o Y, > J3 GO s * I (-1 II ~ 35 * II ii 5 i> OS o || CO io 1 en B* O a e» II II a 02 1-1 II II §5 if " D t5 N Q 1 H M R 11 -=» U9 IO CO CO 1 « I O 1 o II « o II s 1 II II a II e3 II II ^ 1 ■**< ^ Of 1 l-H i „ 11 CO | 6 OS II « ii s "4- § C5 M II eo « 8 " a ii CO 13 II eft W 1 ^ CO CO C& _ i> M II II 1 j OS rt CO 15 II 3 "^ O 1 to 1 1 d B t- II II 8 ■2 II II 'So CO 11 , II . o 13 M CO O 1 1 ' to H ^H M CO TJ< O CO e- CO en o rH M T^ »H H CO THE PERIODIC LAW. MENDELEEFF'S TABLE II. 175 I. n. III. IV. V. VI. RjO I. Li= 7 K 39 Eb 85 Cs 133 — — — — EO n. Be= 9 Ca 40 Sr 87 Ba 137 — — — — B3O3 in. B = 11 Sc 44 T 89 La 138 Yb 173 — — ROj IV. (H,C) C =12 Ti 48 Zr 90 Co 142 — — Th231 K a 6 V. hydrogen compounds decreases. Anhydrides, or Acidic Oxides. — An oxide which, like iodine pentoxide, forms an acid when dissolved in water, or which forms salts by treatment with basic hydroxides, is called an anhydride or acidic oxide. The oxides of the base-forming elements form bases when dissolved in water, and they are, therefore, called basic oxides. As examples of acidic oxides or anhydrides, there may be mentioned besides iodic anhydride, sulphuric anhydride, SO s ; sul- phurous anhydride, 'S0 2 ; phosphoric anhydride, P 2 6 ; carbonic anhydride, C0 2 . When dissolved in water these oxides are converted into acids as represented in these equations : SO s + H 2 = H 2 S0 4 ; S0 2 + H 2 = H 2 SO,; P 2 6 + H 2 = 2HP0 3 ; C0 2 + H 2 = H 2 C0 3 . Silicic anhydride, Si0 2 , is an example of an acidic oxide which does not dissolve in water, but which does form salts when treated with basic hydroxides : Si0 2 + 2KOH = K 2 SiO s + H 2 0. As examples of basic oxides or oxides which when treated ■with water yield bases, the following may be taken : Calcium oxide, CaO; potassium oxide, K 2 0; barium oxide, BaO. PERIODIC ACID. 195 As has already been shown, when treated with water these are respectively converted into calcium hydroxide, Ca(OH) 2 ; potas'sium hydroxide, KOH; and barium hy- droxide, Ba(OH) 2 . There are, however, many basic oxides *which do not dissolve in water, but which, nevertheless, have the power to neutralize acids and form salts. This is true, for example, of aluminium oxide, A 2 10 3 ; lead oxide, PbO; manganous oxide, MnO; cupric oxide, CuO, etc. The action of such oxides upon acids takes place as represented below: A1 2 3 + 3H 2 S0 4 = A1 2 (S0 4 ) 3 + 3H 2 0; PbO + 2HN0 3 = Pb(N0 3 ) 2 + H 2 0; MnO + 2HC1 = MnCl 2 + H 2 0; CuO + H 2 S0 4 = CuS0 4 + H a O. Periodic Acid, H 5 I0 6 . — This acid is analogous to per- chloric acid. Its salts are formed by oxidation of iodates or by heating iodates, just as perchlorates are formed by heating chlorates. The simplest way to make a periodate is to pass chlorine into a solution containing sodium hydroxide and sodium iodate, when a reaction takes place which is at least partly represented by the following equa- tion: 1 NaIO s + 3NaOH + Cl 2 = Na 2 H 3 I0 6 + 2NaCl. The salt Na 2 H 3 I0 6 is difficultly soluble in water, and therefore separates from the solution. From the sodium salt the corresponding silver salt, Ag 2 H 3 I0 6 , can be obtained, and when this silver salt is treated with nitric acid it is converted into the simpler salt, AgI0 4 , which is evidently derived from the simpler acid, HI0 4 : 2Ag 2 H 3 I0 6 + 2HNO, = 2AgN0 3 + 4H 2 + 2AgI0 4 . The acid when separated from its solutions is a crystallized solid which has the composition H 6 I0 6 . When heated it 196 COLLEGE CHEMISTRY. undergoes decomposition, losing water and oxygen, and yielding iodic acid. It cannot, however, be converted into a compound of the composition HI0 4 , for the loss of water is always accompanied by a loss of oxygen. Like iodic acid, periodic acid is a good oxidizing agent in conse- quence of the ease with which it gives up its oxygen. Compounds of Iodine with Chlorine. — When chlorine is passed over dry crystallized iodine it is absorbed, and a compound of the formula IC1 is formed. This is a thick reddish-brown, very volatile liquid. Under proper condi- tions it solidifies in crystals, of which two kinds with different melting-points, 14° and 27°, have been described. Iodine chloride is decomposed by water, the products being iodic acid, hydrochloric acid, and free iodine. If the passage of chlorine over iodine is continued beyond the point required for the formation of the simple compound 101, the trichloride 101, is formed. This is a crystallized compound of a yellow color. When heated it breaks down into chlorine and iodine monochloride. When treated with water it is partly dissolved without decomposition, but it is partly decomposed, yielding iodic acid, iodine monochloride, and hydrochloric acid. Compound of Iodine with Bromine. — There is only one compound of iodine and bromine known, and that is the one having the formula IBr. It is a crystallized compound which is formed by direct combination of the two elements. It is decomposed by heat and by water. Fluorine, F (At. Wt. 19). Occurrence. — This element occurs in large quantity in nature, and is widely distributed, but it is always in com- bination with other elements. It is found chiefly in com- bination with calcium, as fluor-spar or calcium fluoride, CaF 2 , and in combination with sodium and aluminium, as cryolite, a mineral which occurs abundantly in Greenland and has the composition represented by the formula FLUORINE. 197 N"a 3 AlF 6 or A1F 3 . 3K"aF. It is called fluorine from the fact that it occurs in fluor-spar, which in turn receives its name for the reason that it melts when heated and is therefore used as a flux in heating chemical substances together (from fluo, I flow). On account of the remarkable power of combination of fluorine with other elements, all attempts to prepare it in the free condition failed until a few years ago, when its isolation was effected by passing an electric current through a solution of acid potassium fluoride, KHF 2 , in dry hydrofluoric acid contained in a platinum vessel. It has since been shown that a copper vessel may be used. Properties. — Fluorine is a light greenish-yellow gas, that has recently been converted into a liquid at a very low temperature. It acts upon almost all substances. Thus, it decomposes water, yielding ozone and hydrofluoric acid; it combines directly with hydrogen at the ordinary tem- perature; and with sulphur, phosphorus, iron, etc., with evolution of light and heat. It does not, however, act upon platinum. Owing to its active properties it is of course a difficult matter to isolate and preserve it. Hydrofluoric Acid, HF. — Hydrofluoric acid is made by treating a fluoride with sulphuric acid. Thus, when calcium fluoride or fluor-spar is used, this reaction takes place : CaF 2 + H 2 SO, = OaSO, + 2HF. The reaction must be performed in vessels of platinum or lead, as glass is disintegrated by the acid. In perfectly pure anhydrous condition it can be obtained by heating the pure dry salt KHF 2 , known as acid potassium fluoride. It is a liquid which boils at 19.4° and does not solidify even at a very low temperature. The pure dry acid in the liquid form does not act upon glass. It does not dissolve the acid-forming elements, but does dissolve most of the base-forming elements with evolution of hydrogen and formation of fluorides. The gas acts upon the skin, I9 8 COLLEGE CHEMISTRY. causing swellings and violent pains. Inhaled it is poison- ous. To preserve it, vessels of platinum or caoutchouc must be used. In the moist condition it attacks glass, converting the silicon into the fluoride, SiF 4 , and the metals into their fluorides. A silicate of the formula CaSiO s would undergo the changes represented in the fol- lowing equation: fbaSiO s + 6HF = CaF,, + SiF 4 + 3H a O.~T Silicon fluoride is a gas, and calcium fluoride is soluble in acids. Thus calcium silicate, which is insoluble in water, is so changed by hydrofluoric acid as to be rendered solu- ble. In a similar way glass, which is a compound resem- bling calcium silicate, is rendered soluble, or is, as we commonly say, dissolved, by hydrofluoric acid. When an aqueous solution of hydrofluoric acid is boiled it passes over at 120°, and the distillate contains 36 to 38 per cent of the acid. Hydrofluoric acid is used for the purpose of etching glass, particularly for marking scales on thermometers and other graduated glass instruments. The glass is covered with a thin layer of wax or paraffin and, at the places where the etching is wanted, marks are made through the paraffin, so that the glass is exposed. Those parts of the glass which are covered are not acted upon by the hydro- fluoric acid, while those parts which are not covered are corroded and, when the paraffin is removed, permanent marks are found corresponding to those made through the paraffin. A solution of hydrofluoric acid in water is manufactured and sold in rubber bottles. The specific gravity of hydrofluoric acid gas at about 100° leads to the molecular weight corresponding to the formula HF, fluorine having the atomic weight 19. At about 30° the specific gravity corresponds to the formula H 2 F 2 . At lower temperatures the molecular weight appears to be still greater. MANGANESE. 1 99 Family VII, Geoup A — Manganese. Manganese belongs to the same family as the halogens, and resembles them in some respects, but at the same time it differs from them quite markedly in other respects. It acts in fact in two different ways, and is one of those elements, already referred to, that are both acid-forming and base-forming. Some of its compounds with hydrogen and oxygen are distinctly acid, others are distinctly basic. So far as it acts like the members of the chlorine family a brief reference to it here is desirable. On the other hand, it will be dealt with chiefly in connection with those base- forming elements which it most resembles, as, for ex- ample, iron. Manganese occurs in nature principally in the form of pyrolusite or manganese dioxide, Mn0 2 , also known as the black oxide of manganese. It forms with oxygen com- pounds of the following formulas : MnO, Mn 2 3 , Mn 3 4 , Mn0 2 , and Mn 2 0,. When a compound of manganese is subjected to the influence of powerful oxidizing agents in the presence of an alkali it is converted into a salt of manganic acid, H 2 Mn0 4 , which in its composition resem- bles sulphuric acid. If the salt of manganic acid thus obtained is dissolved in water it undergoes partial decom- position, which is complete if the solution is boiled, or if carbon dioxide is passed through it. The change consists in the transformation of manganic acid into permanganic acid, HMn0 4 : 3H 2 Mn0 4 = 2HMn0 4 + Mn0 2 + 2H 2 0; or 3K 2 Mn0 4 + 2H 2 = 2KMn0 4 + Mn0 2 + 4KOH. Permanganic acid, HMn0 4 , is a compound which in many respects resembles perchloric acid. It can be obtained in water solution by decomposition of certain of its salts, but like perchloric acid it is easily decomposed. In consequence of the ease with which it gives up oxygen it is a good oxidizing agent, and is extensively used in the 200 COLLEGE CHEMISTRY. laboratory in this capacity. It is employed in the form of the potassium salt, potassium permanganate, KMn0 4 , which will receive special attention under the head of Manganese" Compounds. In order, however, to make clear the difference in conduct between perchloric and permanganic acids a Jew characteristic facts will be men- tioned here. The conduct of permanganic acid and of potassium permanganate will be understood, if it is borne in mind that in the presence of substances of strongly acid character manganese tends to act as a base-forming ele- ment, and in this capacity to form salts with the acids. Thus in presence of sulphuric acid potassium permanganate forms potassium sulphate, manganous sulphate, and oxygen, if there is anything present which has the power to take up oxygen. In the salts in which it plays the part of a metal manganese is generally bivalent. With hydro- chloric acid, as we have already seen in studying the action of hydrochloric acid upon manganese dioxide, it forms the chloride Mn01 2 . When now potassium permanganate is treated with hydrochloric acid it is decomposed according to the following equation : 2KMn0 4 + 6HC1 = 2KC1 + 2MnCl 2 + 3H a O + 50. Similarly, with sulphuric acid manganous sulphate is formed, thus: 2KMn0 4 + 3H 2 S0 4 = K a S0 4 + 2MnS0 4 + 3H a + 50. Such reactions do not take place with perchloric acid, as chlorine is entirely lacking in the power to enter into acids in the place of the hydrogen and form salts. Manganese forms some other acids besides permanganic acid, but they exhibit little or no analogy with compounds of chlorine, and their study will therefore be postponed until manganese is taken up. The point of chief interest to be noted here is that this element is unmistakably like chlorine in its highest oxygen compounds, but entirely EXPERIMENTS IVITH BROMINE, ETC. 201 different from it in most of its compounds. The com- pound manganese heptoxide, Mn 2 0, , stands in the relation of an anhydride to permanganic acid. In water solution it passes over into the acid : Mn 2 0, + H 2 = 2HMn0 4 . It is formed hy treating potassium permanganate with the most concentrated sulphuric acid : 2KMn0 4 + H 2 S0 4 = K 2 S0 4 + Mn 2 0, + H 2 0. It is extremely unstable, giving up oxygen readily. In contact with organic substances or other substances which have the power to take up oxygen it decomposes so rapidly as frequently to lead to explosions. It is plainly analogous to the oxide of chlorine of the formula CLO,. EXPEEIMENTS. Preparation of Bromine. Experiment 81. — Mix together 3.5 grams potassium bromide and 7 grams manganese dioxide. Put the mixture into a 500 cc. . flask ; connect with a condenser (see Fig. 28). Mix 15 cc. con- centrated sulphuric acid and 90 cc. water. After cooling pour the liquid on the mixture in the flask. Gently heat, when bromine will be given off in the form of vapor. A part of this will con- dense and collect in the receiver. Perform this experiment under a hood with a good draught. Htdrobromic Acid. ^ Experiment 82.— In a small porcelain evaporating-dish put a few crystals of potassium bromide. Pour on them a few drops of concentrated sulphuric acid. The white fumes of hydrobromic acid and the reddish-brown vapor of bromine are noticed. Treat a few crystals of potassium or sodium chloride in the same way. What difference is there between the two cases ? The preparation of hydrobromic acid may be shown in the lec- ture-room as follows : 202 COLLEGE CHEMISTRY. Experiment 83. — Arrange an apparatus as shown in Fig. 40. In the flask put 1 part red phosphorus and 2 parts water. Let Fia. 40. 10 parts bromine gradually drop into the flask from the glass- stoppered funnel. Pass the gas through a U-tube loosely packed with asbestos containing moistened red phosphorus in order to free the hydrobromic acid from bromine, which to some extent passes over with it. Collect some of the gas in water, and examine the solution. How does the gas act when allowed to escape in the air ? Fill a cylinder with the gas in the same way as was done with hydrochloric acid, and fill another with chlorine. While covered with glass plates bring their mouths together. Then with- . draw the plates. "What change is observed ? "What is this due to ? ^ Experiment 34. — To a dilute solution of sodium hydroxide add bromine water made by shaking up a little liquid bromine in a bottle with water. What change takes place ? Add sulphuric acid until the liquid shows an acid reaction. What takes place ? The changes here referred to are perfectly analogous to those which would take place if chlorine were used instead of bromine. Shake a solution containing a little free bromine with ether ; with chlor- oform ; with carbon disulphide. What changes do you observe ? Iodine. -^ Experiment 85. — Mix about 2 grams of sodium or potassium iodide and 4 grams manganese dioxide. Treat with a little con- centrated sulphuric acid in a one- to two-litre flask. Heat gently v." EXPERIMENTS WITH IODINE, ETC. 203 on a sand-bath. Gradually the vessel will be filled with the beau- tiful colored vapor of iodine. In the upper parts of the flask some of the iodine will be deposited in the form of crystals of a grayish-black color. \. Experiment 86.— Make solutions of iodine in water, in alcohol, and in a water solution of potassium iodide. Use small quantities in test-tubes. Experiment 87. — Dissolve a piece of potassium iodide the size of a small pea in about 100 cc. water in a stoppered cylinder. Add enough carbon disulphide to make a layer about an inch thick at the bottom of the cylinder. Shake the two liquids to- gether. Does the carbon disulphide become colored ? Add a drop of chlorine water and shake again. "What difference do you observe in the two cases ? Explain this. Try the same experi- ment, using chloroform instead of carbon disulphide. Iodine can be detected by Means of its Action upon Starch-paste. Experiment 88. — Make some starch-paste by covering a few grains of starch in a porcelain evaporating dish with cold water, grinding this to a paste, and pouring 200-300 cc. boiling-hot water on it. After cooling add a little of this paste to a dilute water solution of iodine in a solution of potassium iodide in water. The solution will turn blue. Now add a little of the paste to a dilute water solution of potassium iodide alone. Is there any change ? Add a drop or two of a solution of chlorine in water. Why the difference ? Will not chlorine water alone act this way toward starch-paste ? Action of Sulphuric Acid upon Potassium Iodide. ■*» Experiment 89. — Bring a piece of potassium iodide the size of a pea in a dry test-tube ; add one drop of water and three or four drops of concentrated sulphuric acid ; the salt becomes brown ; heat gently ; violet-colored vapor escapes, and with it a gas with an odor like that of rotten eggs. At the same time a yellow coating appears on the inside of the tube above the acid. Add five or six drops more of the acid and continue to heat gently. The bad odor first noticed disappears gradually, and another, quite differ- ent odor, irritating to the throat, is now perceptible. This is sul- phur dioxide, SO a . 204 COLLEGE CHEMISTRY. X Iodic Acid. Experiment 90.— Pass chlorine into a test-tube containing iodine in suspension in water ; or add chlorine water. "What be- comes of the iodine ? \, Experiment 91. — Add chlorine water to a dilute solution of potassium iodide, and note the successive changes. \ Experiment 92.— Dissolve iodine in caustie soda. Add an acid to the solution. Explain the changes. Hydrofluoric Acid. ^ Experiment 93.— In a lead or platinum vessel put a few grams (5-6) of powdered fluor-spar and pour on it enough concentrated sulphuric acid to make a thick paste. Cover the surface of a piece of glass with a thin layer of wax or paraffin, and through this scratch some letters or figures, so as to leave the glass ex- posed where the scratches are made. Put the glass over the ves- sel containing the fluor-spar, and let it stand for some hours. Take off the glass, scrape off the coating, and the figures which were marked through the wax or paraffin will be found etched on the glass. CHAPTER XIII. THE ELEMENTS OF FAMILY VI, GROUP B: SULPHUR.— SELENIUM.— TELLURIUM. Introductory. — The elements of this group bear to oxygen a delation somewhat similar to that which the ele- ments of Group B, Family VII, bear to fluorine. The three members sulphur, selenium, and tellurium resemble one another fully as strikingly as chlorine, bromine, and iodine do. Their compounds bear a general resemblance to those of oxygen, and yet "they form very characteristic compounds with oxygen, while oxygen forms no analogous compounds with any of them. Just as iodine forms a compound with fluorine of the formula IF 5 , but fluorine does not form with iodine a compound PI 5 , so sulphur, selenium, and tellurium form with oxygen the compounds S0 3 , SeO,, and TeQ 3 , while oxygen does not form analogous compounds with these other elements. The valence of the elements of this group towards hydrogen is 2, as shown in the compounds H 2 0, H 2 S, H 2 Se, and H 2 Te. Of oxygen and sulphur there are other hydrogen com- pounds, as hydrogen dioxide, H s 2 , and an analogous compound of sulphur, but there is much uncertainty in regard to the constitution of the latter compound, and practically nothing is known in regard to the valence of sulphur in it. In general the hydrogen valence is con- stant. Towards the members of the chlorine group the valence varies from 2 to 6 Oxygen never exhibits a higher valence than 2 towards chlorine and its analogues. The compound OCl 2 illustrates the .bivalence of oxygen towards chlorine. Sulphur forms with chlorine the com- pounds S a Cl a and SC1 3 , which are analogous to the hy- 205 206 COLLEGE CHEMISTRY. drogeu compounds H 2 S 2 and H 2 S, and in both of them the sulphur is probably bivalent. It also forms the compound SC1 4 , in which it is quadrivalent. With iodine it is said to form the unstable compounds S 2 I 2 and SI 6 . With fluorine it forms the stable fluoride, SF 6 . Selenium and tellurium form similar compounds, and, in general, the stability of the compounds of the members of the sulphur group with the members of the chlorine group increases in the order sulphur, selenium, tellurium. Towards oxygen the three elements of the sulphur group are quadrivalent and sexivalent, as seen in the com- pounds S0 2 , Se0 2 , Te0 2 , and SO,, Se0 3 , and TeO,. Of the three elements of this group sulphur occurs in greatest abundance in nature, selenium next in order, and finally tellurium. Just as bromine frequently accompanies chlorine, so selenium frequently accompanies sulphur, but it is always present in much smaller quantity than sulphur. Tellurium occurs in very small quantity relatively, and not uncommonly in combination with valuable metals like gold and silver. Large quantities of sulphur are found in the native or uncombined condition. Only extremely small quantities of selenium and tellurium are found native. Sulphuk, S (At. Wt. 32.06). Occurrence. — The principal deposits of native sulphur are found in Sicily, Italy, and Spain. In California also there is a considerable deposit. In general, sulphur is found near dying or extinct volcanoes. Sulphur occurs in nature, further, in the form of the hydrogen compound, hydrogen sulphide, H 2 S, issuing from the earth in volcanic regions, and in solution in some natural waters, known as "sulphur waters." The oxide S0 2 is likewise found issuing from the earth in volcanic regions. Compounds of sulphur with metallic elements, as with iron, copper, lead, zinc, are very abundant. Such compounds are iron pyrites, PeS 2 ; copper pyrites, CuFeS 2 ; galenite, PbS; and zinc blende, ZnS. Some sulphates are widely distributed SULPHUR. 207 and occur in large quantities; for example, gypsum or calcium sulphate, CaSO t -f- 2H 2 ; barium sulphate, or heavy spar, BaSO^; lead sulphate, PbS0 4 . Finally, sul- phur occurs in a few animal and vegetable products in combination with carbon, hydrogen, and, generally, with nitrogen. Extraction of Sulphur from its Ores. — By far the largest quantity of sulphur found in the market is taken from the mines in Sicily. Of these mines there are between 250 and 300. When taken from the mines it is mixed with many earthy substances, from which it must be separated. In Sicily the separation is generally effected by piling the ore so as to leave passages for air, covering the piles with earthy matter to prevent free access of air and then setting fire to them. A part of the sulphur burns, and this causes the rest of it to melt. The molten sulphur runs down to the bottom of the pile, and by a proper arrangement it is drawn off from time to time. If the pile of ore were not covered up, and the air were allowed to enter freely, the sulphur would burn up, and be converted into the gas sulphur dioxide. The " crude brimstone " obtained in the manner described is afterwards refined by distillation, and it is this refined or distilled sulphur which is met with in the market under the names "roll brimstone" and "flowers of sulphur." The distillation is carried on in retorts made of earthen- ware, and these are connected with large chambers of brick-work. When the vapor of sulphur first comes into the conden sing-chamber it is suddenly cooled, and henoe deposited in the form of a fine powder. This is called " flowers of sulphur." After the distillation has continued for some time the vapor condenses in the form of a liquid, which collects at the bottom of the chamber. This is drawn off into slightly conical wooden moulds, and takes the form of "roll brimstone" or "stick sulphur." Some sulphur is obtained from iron pyrites by heating in closed vessels. The change which takes place on heat- ing iron pyrites is perfectly analogous to that which takes 2o8 COLLEGE CHEMISTRY. place on heating manganese dioxide in preparing oxygen. A sulphur compound of the formula Fe 3 S 4 and free sulphur are formed in the former case, as the compound of man- ganese and oxygen, Mn 3 4 , and free oxygen are formed in the latter: 3FeS a = Fe 3 S 4 + S a ; 3Mn0 2 = Mn 3 4 + O a . Properties. — Sulphur is a yellow, brittle substance which at — 50° is almost colorless. If heated rapidly it melts at 114.5°, forming a thin, straw-colored liquid. When heated to a higher temperature it becomes darker in color, and at 200° to 250° it is so viscid that the vessel in which it is contained may be turned upside down without danger of its running out. Finally, at 448.4° it boils, and is then converted into a brownish-yellow vapor. When molten sulphur solidifies, or when it is deposited from a solution, it takes the form of crystals. But, strange to say, the crystals formed from molten sulphur are entirely different from those deposited from cold solutions of sulphur. The latter belong to the rhombic system. They are octahedrons with a rhombic base. And this is the form of the sulphur found in nature. The former are honey-yellow needles. An examination of these needles shows that the angles which their faces form with one another are not the same as the angles formed by the faces of the octahedrons, and that they belong to an entirely different system — the monoclinic. The rhombic crystals of sulphur can be made by dissolv- ing "roll brimstone." in carbon disulphide and allowing the solution to stand. When the liquid has sufficiently evaporated, the sulphur will appear in larger or smaller rhombic crystals, according to the conditions. A com- parison of the crystals thus obtained with natural crystals will show that the two have identical or very similar forms. The formation of the needles or monoclinic crystals may be shown by melting a considerable quantity, say a pound SULPHUR. 209 or two, of roll brimstone in a sand or Hessian crucible, and allowing tbe liquid mass to cool slowly. When a thin crust has formed on the surface this should be perforated, and the remainder of the liquid sulphur poured off. The crucible will then be found lined with long, dark-yellow, lustrous needles which do not look at all like those obtained from the solution in carbon disulphide. If the monoclinic needles are allowed to lie unmolested they gradually- undergo change spontaneously. They lose their lustre and become lighter in color; and now they consist of minute crystals like those found in nature. They have changed to the rhombic form. It is evident, therefore, that the arrangement of the particles in the monoclinic crystals of sulphur is not a stable one at ordinary temperatures, whereas it has been shown that it is stable at temperatures above 100°. The change from the monoclinic to the rhombic form is accompanied by a considerable evolution of heat. Monoclinic sulphur melts at 120°. Substances which crystallize in two distinct forms are called dimorphous. We shall see that carbon also crystal- lizes in two different forms, and that this kind of relation is not unfrequently met with among chemical compounds. Tbe difference between the two varieties of sulphur sug- gests that observed between the two forms of oxygen. Whether the explanation is the same in the two cases is doubtful. Sulphur can also be obtained in the amorphous, or un- crystallized, condition. If sulphur is heated considerably above its melting-point and then quickly cooled under water it remains for some time soft and dough-like, and while in this condition it is amorphous. If allowed to stand it gradually becomes hard and brittle. When separated from certain compounds which are in solution in water, the sulphur is in a very finely-divided condition, and gives the liquid the appearance of milk. This is seen on adding hydrochloric acid to a solution of sodium thiosulphate, or hyposulphite, as it is generally caL' zio COLLEGE CHEMISTRY. On treating certain varieties of sulphur with carbon disulphide they are found to dissolve completely. This is true, for example, of the natural crystals and of those made artificially by deposition from a solution in carbon disulphide. On the other hand, sulphur in the form of " flowers of sulphur" is only partly soluble in the liquid. There are therefore two forms of sulphur to be distin- guished between, the soluble and the insoluble. The cause of the difference between these modifications is not known. " Stick sulphur" is mostly soluble, while in the " flowers of sulphur " there is at times a considerable percentage of the insoluble variety. Sulphur is insoluble in water, and slightly soluble in alcohol and ether. Sulphur is a much less active element chemically than the members of the chlorine group, and also less active than oxygen. Generally speaking, however, it conducts itself like oxygen. It combines directly though not easily with hydrogen, and it combines readily with most metals, forming compounds called sulphides which, so far as their composition is concerned, are analogous to the oxides. Thus when heated together with iron, copper, or lead, combination takes place readily with evolution of heat and light. In its chemical conduct it differs markedly from the members of the chlorine group. When heated to a suffi- ciently high temperature in the air or in oxygen, sulphur forms the compound sulphur dioxide, S0 2 , and, under certain conditions which will be described farther on, this combines with more oxygen to form the trioxide S0 3 . Further, it combines with nearly all the acid-forming ele- ments if heated with them to a sufficiently high tempera- ture. It combines with most other elements less readily than oxygen and it forms less stable compounds. Thus, its compound with hydrogen is decomposed much more readily into its elements than water is; and the sulphides or its compounds with metals are decomposed when they are heated in oxygen, the oxygen displacing the sulphur, much as chlorine displaces bromine; though there is a SULPHUR. 211 difference between the two cases to be found in the fact that sulphur itself unites readily with oxygen, and this facilitates the decomposition of the sulphides by the action of oxygen. When treated with powerful oxidizing agents in the presence of water sulphur is converted into sulphuric acid. The specific gravity of the vapor of sulphur varies with the temperature, and is such as to indicate that at tem- peratures not far above the boiling-point the molecule con- sists of eight atoms, while at the temperature 800° and higher the molecule consists of two atoms. Uses of Sulphur. — Enormous quantities of sulphur are used in the manufacture of sulphuric acid, and of gun- powder. It is also used in the manufacture of fireworks of various kinds. Burning sulphur gives sulphur dioxide, which is extensively employed for bleaching wool, silk, straw, etc. When caoutchouc is thoroughly mixed with sulphur or some sulphur compound it becomes vulcanized. Compounds of Sulphur with Hydrogen. The principal compound of sulphur and hydrogen is analogous in composition to water. It is known as hydrogen sulphide or sulphuretted hydrogen, and has the formula H,S. Besides this there is at least one other compound which contains a larger proportion of sulphur, and probably has the composition H 3 S a . Further, there are reasons for supposing that still more complex com- pounds can exist, but owing to their instability it has been found impossible to isolate them. Hydrogen Sulphide, Sulphuretted Hydrogen, H 2 S. — When hydrogen is passed over highly-heated sulphur the two elements combine to form hydrogen sulphide. The action is, however, quite incomplete and is not to be com- pared with that which takes place when hydrogen and oxygen are heated together. This compound of sulphur and hydrogen occurs in nature in solution in the so-called 212 COLLEGE CHEMISTRY. " sulphur waters," which are met with in many parts of the earth. It is formed by heating organic substances which contain sulphur, just as water is formed by heating organic substances which contain oxygen. It is formed, further, by decomposition of organic substances which contain sulphur, as, for, example, the albumen of eggs. The odor of rotten eggs is partly due to the formation of hydrogen sulphide. It is formed by the action of acids upon sulphides or hydrosulphides, just as water is formed by the action of acids upon oxides or hydroxides. Thus hydrochloric acid and ferrous sulphide, FeS, give ferrous chloride, FeCl 2 , and hydrogen sulphide : FeS + 2HC1 = FeCl 2 + H 2 S; just as ferrous oxide, FeO, and hydrochloric acid give ferrous chloride and water : FeO + 2HC1 = FeCl 2 + H 2 0. In the laboratory, where the gas is extensively used, it is generally prepared from ferrous sulphide, FeS, and dilute sulphuric acid, which are brought together at the ordinary temperature in a flask such as is used in making hydrogen. The reaction is like that which takes place between ferrous sulphide and hydrochloric acid. It is represented by this equation : FeS + H 2 S0 4 = FeS0 4 + H 2 S. Properties. — Hydrogen sulphide is a colorless, trans- parent gas of the specific gravity 1.178. It has an ex- tremely disagreeable odor, somewhat suggestive of that of rotten eggs. It is poisonous,- even small quantities causing headache, vertigo, nausea, and other bad symptoms. It burns with a blue flame, forming water and sulphur dioxide : H a S + 30 = H 2 + S0 2 . If, however, the air has not free access, as when the gas is COMPOUNDS OF SULPHUR IVITH HYDROGEN. 213 burned in a cylinder open at one end, only a part of the sulphur is burned, the rest being deposited upon the walls of the vessel, while the hydrogen burns. The gas is solu- ble in water, about three volumes being taken up at ordinary temperatures. This solution is used to some extent in the laboratory instead of the gas, but, owing to the fact that it undergoes change in consequence of the action of the oxygen of the air, it is not as valuable as it would be if it were more stable. The change consists simply in the oxidation of the hydrogen and the separation of the sulphur. If a bottle containing a solution of hydrogen sulphide is allowed to stand for a few days, par- ticularly if it is opened from time to time, the odor of the gas will disappear and a deposit of sulphur will be noticed in the bottle. The liquid is then nothing but water. When the solution is boiled it loses all the gas contained in it. Hydrogen sulphide is easily decomposed into its ele- ments. It requires a temperature of only a little above 400° to effect direct decomposition. In consequence of this instability it causes a number of changes which the analogous compound water cannot effect. The relations here are similar to those which exist between hydrochloric and hydriodic acids. Hydrochloric acid is very stable, while hydriodic acid breaks down readily into hydrogen and iodine. Therefore hydriodic acid, as we have seen, acts as a reducing agent, while hydrochloric acid does not. So, also, hydrogen sulphide acts as a reducing agent. Thus, if it is passed into concentrated sulphuric acid this reaction takes place : H 2 S0 4 + H a S = 2H,0 + S + S0 2 . The action is to be traced to the decomposition of the hydrogen sulphide into hydrogen and free sulphur, the hydrogen then acting upon the sulphuric acid thus : H,S0 4 + 2H = 2H 3 -f SO s . 2r 4 COLLEGE CHEMISTRY. With hydriodic acid the reduction may go farther, as has been seen; with hydrobromic acid, however, the action takes place practically in the same way as with hydrogen sulphide. Chlorine, bromine, and iodine act upon hydrogen sul- phide, setting the sulphur free and uniting with the hydrogen : H 3 S + CI, = 2HC1 + S. The instability of hydrogen sulphide is further shown by the ease with which it is decomposed by metals with liberation of hydrogen and formation of sulphides. It will be remembered that at high temperatures several metals decompose water, but that at ordinary temperatures only a few decompose it easily. Hydrogen sulphide acts much more readily; a number of metals which do not act upon water even at high temperatures, as silver, gold, and mercury, decompose this gas at ordinary temperatures. Hydrogen sulphide acts upon metallic oxides, converting them into sulphides, as for example : CuO + H 2 S = CuS + H s O. Action of Hydrogen Sulphide upon Solutions of Salts- Use in Chemical Analysis. — Hydrogen sulphide is exten- sively used in every chemical laboratory as a reagent in chemical analysis. In order that its action may be under- stood a few words of explanation are necessary. Sulphur, as we have seen, readily unites with the metallic or base- forming elements, forming with them the sulphides. Some of these are insoluble in water and in dilute acids; others are soluble in dilute acids but insoluble in water and in ammonia; while still others are soluble even in water. This plainly furnishes a basis for the classification of the base-forming elements into three groups for the purpose of analysis. We have : I. Metals whose sulphides are insoluble in water and are not decomposed by dilute acids. This is called the hy- HYDROGEN SULPHIDE. 215 drogen sulphide group. It includes lead, bismuth, silver, mercury, copper, cadmium, gold, platinum, tin, antimony, and arsenic. II. Metals whose sulphides are insoluble in water but are decomposed by dilute acids. They are therefore not precipitated by hydrogen sulphide, but are precipitated by soluble sulphides. As ammonium sulphide is used for the purpose of precipitating these, the group is known as the ammonium sulphide, group. It includes iron, nickel, cobalt, manganese, thallium, zinc, and uranium. Further, the two elements aluminium and chromium are thrown down with the above, not as sulphides but as hydroxides, and they are therefore included in the group. III. Metals whose sulphides are soluble in water. This group includes all the metals not included in the above two. By taking advantage, then, of the properties of the sul- phides of the metals they can be divided into these three groups, and the detection of any particular element is thus facilitated. If hydrogen sulphide is passed through a solution,, and a precipitate formed, we know that this can contain only those metals which belong to the hydrogen sulphide group; and so on. Now, if the precipitate formed with hydrogen sulphide is treated with certain other reagents other changes take place, and by further study it is quite possible, and indeed comparatively simple, to determine which of the members of the group are present. According to the theory of ions (see pp. 90, 161), hy- drogen sulphide is dissociated in aqueous solution, forming the ions HS and H, and also S and HH. So, also, the metallic salts are dissociated. For example, silver nitrate, AgNO, , in aqueous solution gives the ions Ag and N0 3 . Now when hydrogen sulphide is passed into a solution of silver nitrate the reaction that takes place is to be repre- sented thus : 2(Ag + NO] + S + HH = Ag 3 S + 2H + 3^. 216 COLLEGE CHEMISTRY. The silver sulphide is removed from the solution and is obtained as a solid precipitate, while the ions of nitric acid are in the solution, forming, as we say, a dilute solution of nitric acid. This dilute nitric acid does not act upon the silver sulphide, so that the reaction is complete, all the silver ions being carried out of the solution by the sulphur ions provided enough hydrogen sulphide is passed into the solution. When, however, hydrogen sulphide is passed into a solution of ferrous chloride, FeCl 2 , the results ara differ- ent. In this case, we should naturally expect to find that the reaction would take place as represented in the equa- tion: Fe + 201 + 2H + S = FeS + 2H + 2C1. But the ferrous sulphide, FeS, is acted upon by hydros chloric acid in solution, giving hydrogen sulphide and ferrous chloride, so that the change above represented cannot take place. If, however, a soluble sulphide is added to a solution of a salt of iron or of any metal that acts in the same way, the sulphide is precipitated. Thus, if instead of passing hydrogen sulphide, a solution of potassium sulphide is added, reaction takes place thus : Fe + 201 + 2K + S = FeS + 2K + 201. In the resulting solution the hydrogen ions of hydrochloric acid are not present and the ions of potassium chloride which are present do not act upon the ferrous sulphide. Therefore the ferrous sulphide is precipitated. Hydrosulphides. — The action of hydrogen sulphide shows that it belongs to the class of acids. When it acts upon an oxide the corresponding sulphide and water are formed. But just as there are sulphides that are derived from hydrogen sulphide by the substitution of metallic elements for both hydrogen atoms, so there are hydrosul- HYDROGEN PERSULPHIDE. 217 phides that are derived from it by the replacement of only one of the two atoms of hydrogen in the molecule. The sulphides correspond to the oxides, and the hydrosulphides to the hydroxides. Thus the analogous oxygen and sul- phur compounds of potassium are : K 2 K 2 S KOH KSH. "We speak of sulphides and hydrosulphides as salts of hydrogen sulphide. In consequence of this power to form salts in the same way in general that acids do, hydrogen sulphide is sometimes called sulphydric acid, and the salts of the formula MSH, sulphydrates. The name sulphydrate is analogous to hydrate, which, as has been pointed out, is used by some to designate the compounds of the formula MOH or the hydroxides. Between the acid, hydrogen sulphide, and the neutral compound, water, there is no fundamental difference. The difference is simply one of degree. In the present system of chemistry, which is largely an oxygen system, water is regarded as the con- necting link between acids and bases, as was shown on p. 160. But we might with equal right base our defini- tions and conceptions of acids and bases upon the conduct of sulphur compounds, and thus build up a sulphur system. In such a system hydrogen sulphide would be the connecting link between acids and bases. Hydrogen Persulphide, H 2 S 2 (?). — The sulphides of certain metal?, particularly the so-called alkali metals, sodium and potassium, combine with sulphur to form the polysulphides, examples of which are K 2 S 2 , K 2 S,, K 2 S 4 , and K 2 S 5 . When these are decomposed with dilute acids, compounds of hydrogen and sulphur are formed. It has thus far, however, been impossible to determine whether more than one such compound is formed, for the reason that there is no means of deciding whether the substances formed are chemical compounds or mere mixtures of sul- phur and some one compound of sulphur and hydrogen. 2i8 COLLEGE CHEMISTRY. Hydrogen persulphide is a liquid with a very disagreeable odor. Just as hydrogen dioxide decomposes readily into water and oxygen, so hydrogen persulphide decomposes readily into hydrogen sulphide and sulphur. Compounds of Sulphur with Members of the Chlorine Group. — Sulphur combines directly with chlorine and forms the compounds S 2 C1 2 , SC1 2 , and SC1 4 . Of these the first is the most stable. This can be boiled without under- going decomposition. The second, sulphur dichloride, SC1 2 , undergoes decomposition into chlorine and sulphur monochloride at the boiling-point; while sulphur tetra- chloride exists only at low temperatures. All these compounds are decomposed by water, yielding oxygen compounds. With iodine sulphur forms an unstable hexiodide, SI 6 : with fluorine, a very stable hexafluoride, SF 6 . Selenium, Se (At. Wt. 79.1). Occurrence. — Selenium occurs only in small quantity in nature. It was first found in the deposit formed in a sul- phuric-acid chamber (see p. 227), and owed its origin to the presence of small quantities of selenides in the sul- phides used in the operation. It was found to resemble tellurium, and for that reason was called selenium, from crekrjvr], the moon, tellurium receiving its name from tellus, the earth. Selenium occurs in a number of the natural sulphides, such as iron pyrites, copper pyrites, zinc blende, etc. ; and when these are treated in a current of air to decompose them and convert the sulphur into sulphur dioxide, the selenium is also oxidized, and the selenium dioxide thus formed is carried into the flues and other parts of the apparatus. As it is a solid it is easily condensed, and gradually a considerable quantity collects in the flues. This flue-dust is the best material from which to make selenium. Properties. — There are two modifications of selenium, corresponding to those_of sulphur. One is soluble in TELLURIUM. 219 carbon disulphide, the other is not. The soluble form is obtained by reducing selenious acid by means of sulphurous acid, or by means of other reducing agents. The insoluble variety is obtained by melting selenium, then cooling it down suddenly to 210°, and keeping it at this temperature for some time. When it solidifies it is found to be no longer soluble in carbon disulphide. Selenium burns as sulphur does, forming the oxide SeO a , which has the odor of rotten horse-radishes. Tellurium, Te (At. Wt. 127). Occurrence. — Tellurium occurs in some gold ores in the native or uncombined condition, and also in combination with gold, silver, antimony, lead, and other metals. Tel- lurides occur, among other places in the United States, in California and Virginia. The general method of prepar- ing tellurium from its ores is the same as that by which selenium is made. The tellurium is oxidized to tellurious acid, H 2 Te0 3 , which is isolated, and then reduced by means of sulphurous acid. Properties. — Tellurium is silver- white, and crystallizes easily. Heated in the air it burns, forming a thick white cloud of tellurium dioxide, TeO,. Treated with sulphuric acid it is oxidized to tellurious acid, H 2 Te0 3 , the sul- phuric acid being reduced. Nitric acid oxidizes'it likewise to tellurious acid. Melted together with potassium nitrate it is converted into potassium tellurate, K a Te0 4 . EXPERIMENTS. ' Properties of Sulphur. Experiment 94.— Distil about 10 grains roll sulphur from an ordinary glass retort. What changes in color and in condition take place ? Collect the liquid sulphur formed by the condensa- tion of the vapor in a beaker-glass containing cold water. 2 20 COLLEGE CHEMISTRY. Experiment 95. — Treat some powdered roll sulphur with car- bon disulphide and filter. Does it all dissolve ? Try the same ex- periment with flowers of sulphur. Does this all dissolve ? Put the solutions together and allow to evaporate. Examine the crys- tals deposited. Compare them with some natural crystals of sul- phur. See whether one of the crystals will completely dissolve in carbon disulphide. Experiment 96. — In a covered sand or Hessian crucible melt about 25 grams of roll sulphur. Let it cool slowly, and when a thin crust has formed on the surface make a hole through this and pour out the liquid part of the sulphur. What is left? Compare with the crystals formed in the last experiment. Lay the crucible aside, and in the course of a few days again examine the crystals. What changes, if any, have taken place ? Experiment 97. — Add hydrochloric acid to a solution of sodium thiosulphate. What takes place 1 Experiment 98. — In a wide test-tube heat some sulphur to boiling. Introduce into it small pieces of copper-foil or sheet copper. Or hold a narrow piece of sheet copper so that the end just dips into the boiling sulphur. Experiment 99. — Dissolve some sulphur in concentrated caus- tic soda. In what form is the sulphur in the solution ? y i/ Hydrogen Sulphide. Experiment 100.— Arrange an apparatus as shown in Fig. 41. Put a small handful of the sulphide of iron, FeS, in the flask, and pour dilute sulphuric acid upon it. Pass the evolved gas through a little water contained in the wash-cylinder A. Pass some of the gas into water. [What evidence have you that it dissolves?] Collect some by displacement of air. Its specific gravity is 1.178. Set fire to some of^the gas contained in a cyl- inder. In this case the air has not free access to the gas, and the combustion is not complete. The hydrogen burns to form water, while a part of the sulphur is deposited upon the inside walls of the cylinder. If there is free access of air, the sulphur burns to sulphur dioxide, and the hydrogen to water. Make a solution of the gas in water in the usual way. Put some of this in a bottle and sot it aside, and in the course of a few days examine it again. Boil another portion for a time in a test-tube, and note the changes. Pass a little of the gas through concentrated sulphuric acid contained in a test-tube, and note the HYDROGEN SULPHIDE: EXPERIMENTS. coanges. Moisten strips of paper with dilute solutions of lead nitrate, copper sulphate, stannous chloride, antimony chloride, Fig. 41. and mercuric chloride ; and expose these papers in turn to the gas. What changes take place ? Repeat Experiment 90, and see whether one of the gases given off produces similar changes. Experiment 101. — Pass hydrogen sulphide successively through solutions containing a little lead nitrate, cadmium nitrate, and arsenic prepared by dissolving a little white arsenic, or arsenic trioxide, As a Oi , in dilute hydrochloric acid. What action takes place in each case? The formula of lead nitrate is Pb(NO,) a ; that of cadium nitrate, Cd(NO.)a ; and that of the chloride of arsenic in solution is AsCl 3 . The corresponding sulphides are represented by the formulas PbS, CdS, and As a Si. CHAPTER XIV. COMPOUNDS OF SULPHUR, SELENIUM, AND TEL- LURIUM WITH OXYGEN AND WITH OXYGEN AND HYDROGEN. Introductory. — It has already been stated that, when the three elements; of the sulphur group are burned in the air, they are converted into the corresponding dioxides. Under certain conditions it is possible to convert sulphur dioxide and tellurium dioxide into the trioxides S0 3 and Te0 3 , while the corresponding compound of selenium has not been made. A lower oxide of sulphur, S 2 3 , and a higher one of the formula S 2 0, have also been made. The dioxides dissolve in water, and from these solutions salts of the general formula M 2 SO s ,M 2 Se0 3 , and M 2 Te0 3 are obtained. The trioxide of sulphur dissolves in water with great . evolution of heat, forming compounds S(OH) t , OS(OH) 4 , and 2 S(OH) 2 . Most of the salts obtained from this solution are derived from an acid of the formula H 2 S0 4 , and therefore this is generally called sulphuric acid. By treating sulphuric acid with reducing agents of various kinds it can be con- verted successively into other acids containing a smaller proportion of oxygen; and if the reduction is pushed far enough sulphur and hydrogen sulphide are obtained. The limit of reduction is reached in hydrogen sulphide; and, on the other hand, if hydrogen sulphide is oxidized the limit of oxidation is reached in sulphuric acid. Sulphuric Acid, H 2 S0 4 . — Almost all of the salts of sul- phuric acid can be best explained on the assumption that the forms of sulphuric acid are derived from a compound SULPHURIC ACID. 223 S(OH) 6 . This is called normal sulphuric acid. From normal sulphuric acid by loss of water, the compounds H 4 S0 5 , H 2 S0 4 , and S0 3 are formed : S(OH) 6 = OS(OH) 4 + H 2 0; 0S(0H) 4 = 2 S(OH) 2 + H 2 0; 2 S(OH) 2 = O s S + H 2 0. While these compounds are all known, the salts of sul- phuric acid are for the most part derived from the acid containing two hydrogen atoms, viz., 2 S(OH) 2 . In some salts two of the hydrogen atoms in normal sulphuric acid are replaced by metals, the others remaining. Salts of the general formula S(OH) 4 (OM) 2 are thus formed. These are generally represented as containing two molecules of water of crystallization, thus, M 2 S0 4 + 2H 2 0. To decide be- tween the two views is at present impossible. In the first place, the question as to the nature of water of crystalliza- tion must be answered before it can be said whether there is any conflict between the two views. For the science as well as for the art of chemistry sul- phuric acid is of fundamental importance. It is used daily in every chemical laboratory and in every chemical factory, and in some of the most important branches of chemical industry enormous quantities of it are used. In consequence of the large demand for the acid the process used in preparing it has been studied with great care, and it has reached a high state of perfection. As it furnishes an excellent example of the applications of the facts of science to the building up of an industry, it will be studied with some degree of fulness. Sulphuric acid has been known for a long time. It was made in the eighteenth century by heating calcined iron vitriol (ferrous sulphate, FeS0 4 ) with sand, and was hence called oil of vitriol, a name which still survives. It was also prepared by treating sulphur with saltpetre (potassium nitrate, KN0 3 ). The acid was first prepared on the large scale in England with the use of comparatively large cham- 224 COLLEGE CHEMISTRY. bers lined with lead. It was known as English sulphuric acid, and is still called by this name. Sulphuric acid occurs in nature in the form of salts or sulphates, such as calcium sulphate or gypsum, barium sulphate or heayy spar, and others. Although these salts occur in large quantity, the acid is not obtained from them, as there is no economical way of substituting hydrogen for the metal. The preparation of the acid by the substitution of hydrogen for the calcium in calcium sulphate, 0aS0 4 , suggests itself, but this cannot be easily effected by any substance available in quantity. Hydro- chloric and nitric acids are made from their salts which occur in nature, the former, as we have seen, from sodium chloride, NaCl, the latter from saltpetre or potassium nitrate, KN0 3 , by treating with sulphuric acid. There is, however, no acid that acts upon the sulphates as sul- phuric acid acts upon chlorides, nitrates, and many other salts. The principal process that has up to recently been employed in the manufacture of sulphuric acid is based upon the two fundamental facts that (1) when sulphur is burned it is converted into sulphur dioxide, S0 2 ; and (2) when sulphur dioxide is treated with an oxidizing agent in the presence of water it is converted into sulphuric acid: so 2 + H 2 + = H 2 S0 4 ; or so 2 + H 2 = H 2 SO s ; and H 2 SO, + = H 2 S0 4 . The chief difficulty is, of course, experienced in effecting the oxidation of the sulphurous acid. It is accomplished by introducing the gas, sulphur dioxide, into large cham- bers together with compounds of nitrogen and oxygen, and steam. That which plays the chief part in effecting the transformation is a mixture of the two oxides NO and N0 2 . This acts like the trioxide, N 2 3 , and it may be represented by this formula. Instead of starting with the trioxide, nitric acid is used in the manufacture of sul- SULPHURIC ACID. 225 phuric acid. This at first reacts with sulphur dioxide and steam, as represented in the equation 2HN0 3 + 2S0 2 + H 2 = 2H 2 S0 4 + N 2 3 . After this the main reactions are (1) the formation of a compound of the formula S0 2 < ^^ called nitrosyl-sul- phuric acid; and (2) the decomposition of the nitrosyl- sulphuric acid by water. These reactions are represented in the two equations following : 2S0 2 + N 2 0, + 2 + H 2 = 2S0 2 (OH)(N0 2 ); 2SO *80 ' = °' S \^ ^S0 2 + H,0. uii uv \ 0H H0 / SULPHUROUS ACID. 233 Or the formula of the acid may be written thus : 2 S-OH 2 S-OH. Sulphurous Acid, H 2 S0 3 .— While no acid of the formula H 2 SO s is known in the free condition, a large number of salts which are derived from this acid are known. They are made by treating a water solution of sulphur dioxide with bases, and therefore it is believed that the solution contains the acid which is formed by the action of sulphur dioxide on water, thus: S0 2 + H 2 = H 2 S0 3 . It is, however, so unstable that it breaks down into the dioxide and water at every attempt to isolate it. The dioxide, as has been stated and as will be shown more fully, is formed by the burning of sulphur and by the reduction of sulphuric acid. The acid forms a number of unstable hydrates, apparently of complicated composition. Owing to their great instability, however, the investigation of these substances is extremely difficult and unsatisfactory. Sulphurous acid takes up oxygen readily and is thus transformed into sulphuric acid. It is only necessary to allow a solution to stand for a time to find that the odor of the gas disappears and that sulphuric acid is then present in the solution. Sulphurous acid is frequently used in the laboratory as a reducing agent. The sulphites take up oxygen to form sulphates, and they also take up sulphur to form thiosulphates. The two reactions appear to be perfectly analogous : Na 2 S0 3 + = Na 2 S0 4 ; ]STa 2 S0 3 + S = Na 2 S 2 3 (or N 2 S0 3 S). Sulphurous acid forms two series of salts, the normal sulphites of the general formula M 2 S0 3 , and the acid sul- 234 COLLEGE CHEMISTRY. phites of the general formula MHSO s . These are unstable, though much more stable than the acid itself. When treated with most acids they are decomposed, yielding sulphur dioxide instead of sulphurous acid. The decom- position of sodium sulphite with hydrochloric acid is represented by the equation Na 2 SO, + 2HC1 = 2NaCl + H 2 + S0 2 ; with sulphuric acid thus : Na 2 S0 3 + H 2 S0 4 = Na 2 S0 4 + H 2 + S0 2 . Hyposulphurous Acid, H 2 S 2 4 . — This acid is also called hydrosulphurous acid, but the name hyposulphurous acid is more in accordance with the nomenclature adopted for the other acids, and is now preferred. But little is known of the compound. It is formed by reduction of a salt of sulphurous acid by zinc: Zn + 2S0 2 = ZnS 2 4 . The free acid cannot be obtained from this salt. Thiosulphuric Acid, H 2 S 2 3 , — This acid was formerly, and is still by many, called hyposulphurous acid. Its formation, or the formation of its salts by the addition of sulphur to the sulphites, has been mentioned, and the analogy between this reaction and that of the formation of sulphates by the addition of oxygen to sulphites has been commented upon. It may be regarded as sulphuric acid in which one atom of sulphur has been substituted for one atom of oxygen, and hence the name thiosulphuric acid is appropriate, whereas the name hyposulphurous acid suggests at once a compound similar to sulphurous acid, but containing less oxygen, and is therefore inappropriate. The acid itself is very unstable, breaking down into sulphur dioxide, sulphur, and water. By acids its salts are decomposed in a similar way with evolution of sulphur dioxide and separation of sulphur, which appears sus- SULPHUR DIOXIDE. 235 pended in the liquid in a very fine state of division. With hydrochloric acid the decomposition takes place thus: Na 2 S 2 O s + 2HC1 =.2NaCl + SO, + S + H 2 0. Other Acids of Sulphur — Of the other acids of sulphur but little need be said here. They form a series the members of which are closely related to one another as shown below: Dithionic acid H a S 2 6 Trithionic acid H 2 S,0 6 Tetrathionic acid H 2 S 4 O s Pentathionic acid H 2 S 5 6 Persulphuric Acid, HS0 4 , is formed by dissolving the oxide S 2 0, in water, the oxide itself being formed by sub- jecting a mixture of sulphur dioxide and oxygen to the silent discharge in an ozone tube (see p. 116). The potas- sium salt of persulphuric acid is easily obtained by sub- jecting a saturated solution of acid potassium sulphate to electrolysis. Caro's Reagent.— When potassium persulphate is added to moderately dilute sulphuric acid a solution is obtained that shows remarkable oxidizing powers. It has recently come into use in laboratories under the name of Caro's reagent. Compounds of Sulphur with Oxygen. — Sulphur, as has been repeatedly stated, combines with oxygen in two pro- portions, forming the oxides S0 2 and S0 3 , or sulphur dioxide and sulphur trioxide. Besides these two it also forms a sesquioxide, S 2 3 , and a heptoxide, S 2 0, , but com- paratively little is known in regard to the last two. The one best known is sulphur dioxide. t«*~ Sulphur Dioxide, S0 2 . — This, as has been seen, is "^formed when sulphur is burned in the air or in oxygen; and it is also easily formed by reduction of the higher oxides and acids of sulphur. Owing to the fact that with 236 COLLEGE CHEMISTRY. water it forms sulphurous acid, it is frequently called sul- phurous anhydride. The methods for making it were referred to under sulphuric acid, and the reactions involved were sufficiently discussed. It need only be said here that in the laboratory the methods most commonly employed are: (1) Heating sulphuric acid with copper; (2) heating the acid with carbon (charcoal) ; (3) heating the acid with sulphur; and (4) allowing concentrated sulphuric acid to drop into a 40 per cent solution of acid sodium sulphite: H 2 S0 4 + HNaS0 3 = HNaSO, + S0 2 + H 2 0. Sulphur dioxide is a colorless, transparent gas, which has a pungent, suffocating odor, familiar as the odor of burning sulphur matches. It is poisonous, causing death when inhaled in any quantity, and giving rise to bad symptoms even in comparatively small quantities. It does not readily give up its oxygen, so that burning bodies are extinguished when introduced into it. It acts something like water in this respect. It is more than twice as heavy as air, its specific gravity being 2.26. When sulphur is burned in oxygen gas the sulphur dioxide formed occupies the same volume as the oxygen used up, so that there is no change in the volume. This can be shown by the experiment here described. In a bent glass tube, of the form shown in Fig. 43, there is placed a piece of sulphur, and the tube is then half filled with pure oxygen over mercury. Gn now Fio. 43. heating the tube at the part where the sulphur is, this burns and is converted into sulphur dioxide. After the tube has cooled down to the ordinary temperature the gas is found to occupy the same volume as before. This will be readily understood by the aid of the following con- siderations : In the reaction s +■ 0, = so 2 one molecule of sulphur dioxide is formed for every mole- SULPHUR DIOXIDE. 237 cule of oxygen used up. But a molecule of sulphur dioxide in the form of gas occupies the same volume as a molecule of oxygen, so that, as the volume occupied by the sulphur in the experiment is insignificant, there is no change in volume occasioned by the above reaction. Sulphur dioxide dissolves in water, as we have seen, and forms a liquid in which, judging by its conduct, sul- phurous acid is present. The gas is easily liquefied by cold alone. It is only necessary for this purpose to pass the dry gas through a tube surrounded by a freezing mixture of ice and salt. The liquid changes rapidly into gas under ordinary pressure at the ordinary temperature. In this change so much heat is absorbed that a temperature of about — 60° can be pro- duced by means of it; and a portion of the liquid can be solidified. Sulphur dioxide is very stable. If heated to 1200° under pressure, however, it breaks down into sulphur tri- oxide and sulphur : 3S0 2 = 2S0 3 + S. When conducted into solutions of bases or of carbonates the corresponding sulphites are formed : 2K0H + SO, = K 2 S0 3 + H 2 0; K 2 C0 3 + S0 2 = K 2 S0 3 + C0 2 . Under certain conditions, as when the gas is passed into a hot solution of an alkali carbonate, a salt of the general formula M 2 S 2 5 is formed. This bears to "the sulphite the same relation that the pyrosulphate bears to the sulphate : 2KHS0 3 = K 2 S 2 5 + H 2 0; and 2KHS0 4 = K 2 S 2 0, + H 2 0. Sulphur dioxide is used extensively for the purpose of bleaching silk, wool, straw, and basket-ware. In order that it may bleach, however, water must be present, so 238 COLLEGE CHEMISTRY. that it appears that the true bleaching agent in this case is sulphurous acid and not the dioxide. When we consider that sulphur dioxide does not readily take up nor give up oxygen, while sulphurous acid does readily take it up, the necessity of having water present in the bleaching process at once becomes apparent. The bleaching in some cases certainly consists in abstracting oxygen from the colored substances, and thus converting them into colorless products. In other cases it is due to the formation of compounds of sulphurous acid with the dye-stuffs. Sulphur dioxide is not only a bleaching agent like chlorine, but like chlorine it is also a disinfectant. It has to some extent the power to destroy the organisms which cause changes in organic substances. It prevents fermen- tation and is therefore used as a preservative. Its power to destroy the germs of disease, that is, to disinfect, is not as great as is frequently supposed. Much larger quanti- ties are necessary for this purpose than are commmonly used. Sulphur Trioxide, S0 3 . — This compound is made by passing sulphur dioxide and oxygen together over heated platinum in a finely-divided state. It is obtained most readily by heating disulphuric acid, which breaks up easily into sulphur trioxide and ordinary sulphuric acid according to the equation H 2 S 2 0, = H 2 S0 4 + SO s . Similarly, the acid sulphates of the alkali metals yield the corresponding normal sulphates and sulphur trioxide : 2NaHS0 4 = Na 2 S0 4 + S0 3 + H 2 0. It is now manufactured on the large scale by passing sul- phur dioxide and oxygen together over asbestos covered with finely-divided platinum, ferric oxide, and some other substances, and the product thus obtained is converted into sulphuric acid by the action of water, or it is passed into SULPHUR TRIOXIDE. 239 ordinary sulphuric acid for the purpose of making "solid sulphuric acid," which is almost pure disulphuric acid, HAO,. Sulphur trioxide is a white crysfeta^tersolid i#MUh appears to exist in two modifications. The one crystallizes in fine needles like asbestos. When heated it passes directly into gas. The other solid variety melts at 14.8°, forming a liquid which boils at 46°. In contact with the air the oxide gives off thick fumes which are partly due to the great power of the compound to combine with water. , Water acts with violence upon it, the heat evolved in the act being 39,170 cal. It also acts upon substances containing hydrogen and oxygen in much the same way that concentrated sulphuric acid does, charring them by abstracting the hydrogen and oxygen. It acts, however, more violently in this way than sulphuric acid does. With water it forms sulphuric acid, and it is, therefore, called sulphuric anhydride. The reaction involved in passing from sulphur trioxide to sulphuric acid is of a kind which, as we have seen, is frequently met with both with acidic oxides or anhydrides, and with basic or metal- lic oxides; and it is desirable that it should here be studied a little more carefully than it has yet been. What we know is that when sulphur trioxide acts upon water there is a great deal of heat evolved, and compounds of different composition are obtained. The composition of these compounds is represented by the formulas S0 3 + H 2 0, S0 3 + 2H 2 0, and SO, + 3H 2 0. So, too, when calcium oxide or lime, OaO, acts upon water, there is great evolution of heat, and a compound is formed the composi- tion of which is represented by the formula CaO + H 2 0. But in these formulas no attempt is made to give any account of what takes place in the chemical acts referred to. That the water is not present in the compounds as water seems evident, in the first place from the conduct of the compounds, and in the second place from the amount of heat evolved in the act of combination. Now, taking the chemical conduct of the substances into consideration, 24« COLLEGE CHEMISTRY. they appear to contain hydrogen in combination with oxygen, and their conduct becomes comprehensible on the supposition that the group known as hydroxyl (-0-H) is present. This view has been found to be in accordance with a large number of facts, and it is of great assistance in dealing with these facts. The view is distinctly this: When an acidic oxide acts upon water it is converted into a hydroxyl compound which has acid properties, as shown in this equation: II H a 0=S=0 -f- H,0 = S \ H a OH OH OH OH' OH OH According to this, each molecule of water is decomposed and each hydrogen atom in the resulting compound is in combination with oxygen. The same kind of action takes place in the case of some basic oxides, as shown, for ex- ample, in the case of calcium oxide: Ca=0 + H,0 = Ca OH The potassium salt of the acid is formed by melting together tellurium dioxide and potassium carbonate: K a C0 3 + TeO a = K a TeO + CO a . If the salt thus formed is dissolved in water and nitric acid added to the solution, tellurious acid is thrown down : K a Te0 3 + 2HN0 3 = 2KN0 3 + H a Te0 3 . • It is a solid which easily loses water and is thus trans- formed into tellurium dioxide. Telluric Acid, H a Te0 4 . — Salts of this acid are formed by melting tellurious acid with saltpetre and other oxidizing agents. When the solution of the acid is evaporated to crystallization the solid compound deposited has the com- position H„TeO e , and, according to what was learned in studying sulphuric acid, it appears probable that this is normal telluric acid, Te(OH) 6 . When normal telluric acid is heated to a little above 100° it loses water and is trans- formed into the acid HjTeO^ , corresponding to ordinary sulphuric acid, from which most of the tellurates are derived : Te(OH) 6 = TeO a (OH) a + 2H a O. Heated higher, to about 160°, the acid is decomposed into tellurium trioxide and water : TeO,(OH) a = TeO s + H a O. / / 244 COLLEGE CHEMISTRY. Although most of the tellurates are simple salts of the acid H,Te0 4 , others are derived from more complex forms of the acid. One of these is analogous to disulphuric acid. Oxides of Tellurium. — Tellurium monoxide is formed by heating sulphur trioxide and tellurium together in a vacuum. Tellurium dioxide is formed by burning tel- lurium or by oxidizing it with nitric acid; and, further, by the decomposition of tellurious acid by heat. It crystallizes and is but slightly soluble in water. The trioxide, Te0 3 , is formed by heating telluric acid to a high temperature. Its conduct is entirely different from that of sulphur trioxide. "While the latter acts with violence upon water, and readily upon metallic oxides and hydrochloric acid, the former does not act readily upon any of these substances. It is insoluble in hot as well as cold water. Family VI, Group A. Group A, Family VI, includes chromium, molybdenum, tungsten, and uranium. All of these show some resem- blance to the elements of the sulphur group, but they also appear, in entirely different characters, forming compounds of a kind unknown among the derivatives of sulphur and its analogues. The relation which these elements bear to sulphur is much like that which manganese bears to chlorine. The resemblance to sulphur is seen mainly in the formation of acids of the formulas H 2 Cr0 4 , H 2 Mo0 4 , H 2 W0 4 , and H 2 TT0 4 ; and the oxides Cr0 3 , MoO s , W0 3 , and U0 3 . The common salts of chromic acid are derived from dichromic acid, which is analogous to disulphuric acid. They have the general formula M 2 Cr 2 7 . So, too, salts of molybdic acid are known which are derived from the simple form of the acid, H 2 Mo0 4 , and others which are derived from a dimolybdic acid, H 2 Mo 2 0, , and from more complicated forms. Tungsten has a wonderful power of forming complex acids. All of them, however, FAMILY VI, GROUP A. 245 can be referred to the simple form HjWO^. And, finally, nranio acid forms salts which for the most part are derived from diuranic acid, H 2 U 2 0,. All the most important of these compounds will be taken up later. When the acids of chromium, molybdenum, tungsten, and uranium lose oxygen they form compounds that have little or no acid character. The lower oxides of chromium form salts with acids, and these bear a general resemblance to the salts oh aluminium, iron, and manganese. The chromates lose their oxygen quite readily when acids are present with which the chromium can#nter into combina- tion in its capacity as a base-forming element. Thus, when potassium chromate, K 2 Cr0 4 , is treated with hydro- chloric acid in the presence of something which can take up oxygen, decomposition takes place thus : 2K 2 CrO< + 10HC1 = 4K01 + SCrCl, + 5H 2 + 30. With sulphuric acid the action takes place as represented in this equation : 2K 2 Cr0 4 + S^SO, = 2K.S0, + CfctSQJ, + 5H 2 •+- 30. In both these cases the chromium enters into combination as a trivalent base-forming element, taking the place of three atoms of hydrogen in hydrochloric acid in the first case, and of three atoms of hydrogen in the sulphuric acid in the second. Molybdenum and tungsten do not form salts of this character; indeed they seem to be practically devoid of basic properties. Uranium, on the other hand,g forms some curious salts which differ from the simple metallic salts which we commonly have to deal with. 246 COLLEGE CHEMISTRY. EXPERIMENTS. Manufacture of Sulphuric Acid. Experiment 102. — The manufacture of sulphuric acid can he illustrated in the laboratory by means of the apparatus repre- sented in Fig. 44. This consists of a large balloon-flask fitted with a stopper having five openings. By means of tubes it is con- nected with three small flasks. One of these, a, contains water for the purpose of providing a current of steam ; another, c, con- tains copper-foil and concentrated sulphuric acid, which give sul- Fio. 44. » phur dioxide \7hen heated ; and the third, 6, contains copper-foil and dilute nitric acid, which give oxides of nitrogen, mainly nitric oxide, NO. "When the nitric oxide comes in contact with the air it combines with oxygen, forming nitrogen trioxide and H^nitrogen peroxide ; and when steam and snlphur dioxide are ad- mitted to the flask the reactions involved in the manufacture of sulphuric acid take place. By means of a pair of bellows attached at d air is supplied. If air is not forced in, the gases become colorless, owing to complete reduction of the oxides of nitrogen to the form of nitric oxide, NO, which is colorless. If steam is not admitted the walls of the vessel become covered with crystals of nitrosyl-sulphuric acid. This is, however, decomposed by an excess of steam. SULPHUROUS ACID AND SULPHUR DIOXIDE. 247 Experiment 103.— Into a vessel containing ordinary concen- trated sulphuric acid introduce small sticks of wood, pieces of paper, and various other organic substances, and note the result. The charring effect is particularly well shown by adding the acid drop by drop to a concentrated solution of sugar, or to molasses and stirring. Experiment 104. — Sulphuric acid is detected in analysis by add- ing barium chloride to its solution, when insoluble barium sul- phate is formed : H„SO. + Bad, = BaSO. + 2HC1. - >;£. • Other insoluble sulphates are those of swontium and lead ; and calcium sulphate is difficultly soluble. To a dilute solution of sul- phuric acid or of any soluble sulphate, add in test-tubes barium chloride, strontium nitrate, and lead nitrate. Sulphurous Acid and Sulphur Dioxide. ^ Experiment 105. — Put eight or ten pieces of sheet copper, one or two inches long and about half an inch wide, in a 500 cc. flask; pour 15 to 20 cc. concentrated sulphuric acid on it. On heating, sulphur dioxide will be evolved. The moment the gas begins to come off, lower the flame, and keep it at such a height that the evolution is regular and not too active. Pass some of the gas into a bottle containing water. The solution in water is called sul- phurous acid. \ Experiment 106. — The most convenient method for making sulphur dioxide consists in letting moderately concentrated sul- phuric acid drop into a water solution of acid sodium sulphite. Fit a flask with a stopper with two holes. Throflgh one of the holes pass a dropping-funnel and through the other a convenient delivery-tube. Jn the funnel put a cooled mixture of equal vol- umes of concentrated sulphuric acid and water. Tn the flask put a 40-per cent solution of acid sodium sulphite in water. By ♦" turning the stop-cock of the funnel so that the acid drops into the solution there will be a regular evolution of the gas. ~- Experiment 107. — Pass sulphur dioxide into a moderately dilute solution of potassium hydroxide, until the solution is saturated. "What is then contained in the solution ? To a little of it add hydrochloric acid. "What takes place ? ^ Experiment 108. — Try the effect of heating concentrated sul- phuric acid with charcoal, and with sulphur. 248 COLLEGE CHEMISTRY. X Experiment 109. — Collect by displacement of air some of the gas made in Experiment 106. Does it burn 1 *^ or does it support combustion ? Vx> , Experiment 110.— Pass some of the gas through a bent-glass tube surrounded by a freezing mixture of salt and ice. Tubes u Fio. 45. provided with glass stop-cocks are made for such purposes. They generally have the form represented in Fig. 45. If the tube is taken out of the freezing mixture, the liquid sulphur dioxide changes rapidly to gas, if the tube is open. Experiment 111.— Burn a little sulphur in a porcelain crucible under a bell-jar. Place over the crucible on a tripod some flowers. In the atmosphere of sulphur dioxide the flowers will be bleached. Sulphurous Acid is a Reducing Agent. Experiment 112. — To a dilute solution of potassium iodide in a test-tube gradually add chlorine water until the solution be- comes clear and colorless. Now add a solution of sulphurous acid. At first iodine is deposited, but on further addition of sulphurous acid it dissolves again. Explain all the changes. # Sulphur Trioxide. Experiment 113. — Heat a little fuming sulphuric acid gently in a test-tube. What takes place 1 Put a little of the acid (5-10 cc.) in a small dry retort provided with a glass stopper and connect with a dry glass receiver. Heat the retort gently, and keep the receiver cool. By means of a dry glass rod take out some of the substance which collects in the receiver and put it in water. Lay a little of it on a piece of wood and on a piece of paper. Experiment 114. — Prepare finely-divided platinum by moisten- ing some fine asbestos with a solution of platinic chloride and heating to redness in a porcelain crucible. The substance thus obtained is known as platinized asbestos. Now arrange an SULPHUR TRIOX1DE: EXPERIMENTS. 249 apparatus so Uiat both oxygen and sulphur dioxide can be passed together through a tube of hard glass as represented in Fig. 46. Fig. «. First pass the two dried gases together through the empty tube and heat a part of the tube by means of a burner. Is there any evidence of combination ? Now stop the currents of the gases, let the tube cool down, and introduce a small layer of the plat- inized asbestos. Pass the dried gases over the heated asbestos. What ' ak»<» place ? CHAPTEE XV. NITROGEN.— THE AIR.— ARGON, ETC. Nitrogen, N (At. Wt. 14.04). General. — Nitrogen bears to a group of elements rela- tions very similar to those which oxygen bears to the sul- phur group, and fluorine to the chlorine group. There are easily recognized resemblances between it and the members of the group, and yet' there are some marked differences. As has been stated, and as is seen from its position in the periodic system, nitrogen is trivalent towards hydrogen, as shown in the compound NH S , while it is both trivalent and quinquivalent towards oxygen, as appears to be shown in N 2 O s and N 2 6 . The principal hydrogen compound, ammonia, is entirely different in character from those of chlorine and sulphur, for, while these are acid, ammonia has in a marked way the character of a base, acting, however, in a peculiar way upon acids to form salts. The two oxides above referred to are acidic, forming the acids HN0 2 and HN0 3 , which are known as nitrous and nitric acids respectively. Occurrence of Nitrogen. — It was discovered by Lavoisier and Scheele towards the end of the last century that the air consists of two gases, one of which is oxygen, and they showed that when the oxygen is removed the gas which is left has not the power to support combustion nor to sup- port respiration. This gas was first called azote (from a, privative, and Qootikos, life), and this name is still retained in France, the symbol in use in that country being Az, whereas in all others the symbol is N. This is the only case in which there is a difference of usage in respect 250 NITROGEN. 251 to the symbols of the chemical elements in different countries. The name nitrogene was given to it later, from the fact that it is a constituent of nitre or saltpetre, KNO {nitrum, saltpetre, and yiveiv, to produce), and this is the origin of the English name nitrogen. Not only is nitrogen found free in the air, but it is found in combina- tion in a large number of substances in nature. It is found in the nitrates, or salts of nitric acid, particularly as the potassium salt KN0 3 , and the sodium salt NaNO s , which occurs in enormous quantities in Chili, and is therefore known as Chili saltpetre. It is also found in the form of ammonia, which is a compound of nitrogen and hydrogen of the formula NH 3 . Ammonia occurs in small quantity in the air, and is formed under a variety of conditions, which will be taken up later. -Nitrogen occurs, further, in combination in many animal substances. Preparation. — The most convenient way to prepare nitrogen is by burning in a closed vessel something which does not give a gaseous product of combustion; or by pass- ing air over something which has the power to unite with oxygen. The best substance to use for the first purpose is phosphorus, which burns readily and yields a solid product, soluble in water. It is only necessary, therefore, to place a piece of phosphorus in a floating vessel on the surface of water, set fire to it, and immediately place over it a closed bell-jar. As soon as the oxygen is used up the combustion stops, and the vessel then contains the residual nitrogen, and the walls are covered with a thin layer of phosphorus pentoxide, P 2 5 . This is soon converted by the water into phosphoric acid, which dissolves. Another convenient method for preparing nitrogen consists in passing air ovei copper heated in a tube. The copper takes up the oxygen readily, and the nitrogen passes on. Another good method consists in exposing to the air copper turnings partly covered with a solution of ammonia in a vessel so arranged as to allow free access of air while the escape of the gas in the vessel is prevented. This mixture absorbs oxygen slowly at the ordinary temperatures. Nitrogen can also 25? COLLEGE CHEMISTRY. be made from other substances than the air. Thus, when chlorine is passed into a water solution of ammonia this reaction takes place : NH 3 + 3C1 = N + 3HC1; but the hydrochloric acid combines at once with ammonia to form ammonium chloride, NH 4 C1: NH 3 + HC1 = NH 4 C1; so that the only gaseous product is nitrogen. This experi- ment is more or less dangerous, for if all the ammonia should be used up, and the passage of chlorine continued, a compound of nitrogen and chlorine which is extremely explosive is formed. Finally, nitrogen can be made by heating ammonium nitrite, NH 4 N0 2 , either dry or in solution. The hydrogen and oxygen of the compound unite to form water and the nitrogen is set free: NH 4 N0 2 = 2H a O + N 2 . The nitrogen prepared from the air is never pure, as there are always present in the air other substances besides nitrogen and oxygen; and while some of these can be removed without serious difficulty, others cannot be. Properties. — Nitrogen is a colorless, tasteless, inodorous gas. At — 194° under the ordinary atmospheric pressure it is converted into a liquid ; and this liquid solidifies at — 214°. The critical temperature is — 146°, and the critical pressure is 35 atmospheres. A litre of nitrogen under standard conditions weighs 1.2505 grams. Its specific gravity (air — 1) is 0.967. It does not support combustion, nor does it burn. This latter fact is obvious, for,' if nitrogen had the power to combine with oxygen when the temperature of the mixture is elevated, it is plain that this process of combustion would long ago have taken place, leaving one or the other of the two gases and the product of combustion as the constituents of the air. Nitrogen not only does not combine with oxygen readily, >:*& test, to NITROGEN. 233 but it does not combine readily with any other element except at a very high temperature. Just as it does not support combustion, so also it does not support respiration. Animals would die in it, not on account of any active poisonous properties possessed by it, but for lack of oxygen. In the air it serves the useful purpose of diluting the oxygen. If the air consisted only of oxygen, all processes of combustion would certainly be much more active than they now are. "What effect the continued breathing of oxygen would have upon animals it is impossible to say. The Air — The atmosphere of the earth, commonly called the air, consists essentially of the two elements nitrogen and oxygen in the proportion of 79 volumes of nitrogen to 21 volumes of oxygen, or, by weight, of 77 per cent of nitrogen and 23 per cent of oxygen. The weight of a litre of air under standard conditions is 1.293 grams. Wherever air has been collected and analyzed it has been found to have practically the same composition. Nevertheless very accurate analyses have shown that the composition of the air is subject to slight variations. To decide whether the air is a chemical compound or a mechanical mixture requires a careful examination of a number of facts. The evidence may be summed up as follows: (1) If nitrogen and oxygen are mixed together the mix- ture conducts itself in exactly the same way as air. The mixing is not attended by any phenomena indicating chemical action. Generally the chemical combination of two elements is accompanied by an evolution of heat, and whenever a chemical act takes place there is some change in the temperature of the substances. When nitrogen and oxygen are brought together there is no change in the temperature of the gases, and no change of volume. (2) The law of definite proportions is founded upon a very large number of observations, and in all cases in which we have independent evidence that chemical action takes place it is found that the substances combine in exactly the same proportions to form the same product. 254 COLLEGE CHEMISTRY. Variation itt the composition of a chemical compound is not known. The composition of the air varies slightly, according to circumstances, and this fact may be regarded as evidence that the air is not a chemical compound. (3) Air dissolves somewhat in water. If air which is in solution in water is pumped out and analyzed, it is found to have a different composition from that of ordinary air. Instead of containing nearly 4 volumes of nitrogen to 1 of oxygen, it contains only 1.87 volume.s of nitrogen to 1 of oxygen. The proportion of oxygen is much larger in the air which has been dissolved in water than it is in ordinary air. This is due to the fact that oxygen is more soluble in water than nitrogen. Therefore, when air is shaken with water, relatively more oxygen than nitrogen is dissolved. If the gases were in chemical combination we should expect the compound to dissolve as such and without change of composition. The above evidence shows that nitrogen and oxygen are not combined chemically in the air, but that they are simply mixed together. Analysis of Air. — The earliest examinations of Priestley, Lavoisier, and Scheele were made by burning substances in air contained in closed vessels. They concluded that the air is made up of % oxygen and \ nitrogen by bulk. In order to determine the composition of the air to-day, we should proceed as follows: A qualitative examination would easily show the presence of nitrogen and oxygen. If a solution of calcium hydroxide, Ca(OH) 2 , which is known as lime-water, or a solution of barium hydroxide, Ba(OH) 2 , is exposed to the air it becomes turbid, and a precipitate is formed. Neither nitrogen nor oxygen nor an artificially prepared mixture of the two gases can pro- duce this change. It has been shown that the change is due to the presence in the air of a small quantity of the gaseous compound, carbon dioxide, C0 2 . If calcium chloride or phosphorus pentoxide is exposed to the air it soon becomes moist and after a time turns liquid. This effect has been found to be due to the presence of water- ANALYSIS OF AIR. 255 vapor in the air. Nitrogen obtained from the air by passing the latter over copper, which abstracts the oxygen, has been shown to contain a small amount of an extremely inert gas, argon (which see). By other methods which need not be considered here it can be shown that there are many other substances in the air besides those mentioned. Among them are ammonia, hydrogen dioxide, and organic matters of various kinds, including a large variety of germs the presence of which can be detected by the changes which exposure to the air produces in certain liquids, as milk and fruit-juices. Having thus learned what the chief constituents of the air are, the next thing is to determine in what quantities they are present, or to make a quantitative analysis of the air. For this purpose advantage may be taken of the fact that phosphorus when exposed to the air at ordinary tem- peratures combines slowly with the oxygen, leaving the nitrogen, argon, and minute quantities of other gases. If, therefore, a piece of ordinary phosphorus is inserted into a measured volume of air contained in a graduated glass tube over water or mercury, a diminution in volume will take place slowly. If, in the course of a few hours, the volume is again measured, the difference will give the volume of oxygen absorbed, while the gas remaining is nitrogen. Of course, in this case as in all others in which gas volumes are measured, corrections for temperature, pressure, and tension of aqueous vapor must be made. Another method by which the ratio between the nitrogen and oxygen in air can be determined is that which was first employed by Dumas and Boussingault. It consists in passing air over heated copper, collecting and measuring the nitrogen, and weighing the copper oxide. The appa- ratus is arranged as shown in Fig. 47. The copper is contained in the glass tube ab on the combustion furnace. At the ends of this tube are the stop-cocks rr. V is a glass globe provided with a stop- cock u. Before the experiment the air is exhausted from the globe and the tube ab, and the tube then carefully 256 COLLEGE CHEMISTRY. weighed. The tubes B and C and the apparatus A con- tain substances which have the power to absorb the carbon dioxide of the air. The tube ab is now heated and air admitted after passing through G, B, and A. The copper Fio. 47. takes up the oxygen, and the nitrogen enters the globe V. After the globe is full it is weighed, then exhausted and weighed again, and the difference gives the weight of the nitrogen. The tube is also exhausted and weighed, and the difference between this weight and that of the ex- hausted tube before the experiment gives the weight of the oxygen. The most refined method for the analysis of the air is the eudiometric method of Bunsen. This consists in adding some pure hydrogen to a measured volume of air contained in a eudiometer (see p. 74) over mercury, and then exploding the mixture by means of an electric spark. If the conditions are right all the oxygen present will combine with hydrogen, and in consequence of this there will be a corresponding contraction in the volume of the gases. The amount of contraction will be equal to the volume of hydrogen and that of oxygen which have com- bined to form water. But we know from previous experi- ments on these two gases that they combine in the ratio ANALYSIS OF AIR. 257 of one volume of oxygen to two of hydrogen. Conse- quently the volume of oxygen which was present is equal to one-third of the total contraction. Of course it is necessary that there should be enough hydrogen present to combine with all the oxygen. This method is capable of great exactness. The most accurate analyses made by this method by Bunsen and others have shown that in 100 volumes of air there are 20.9 to 21 volumes of oxygen. The estimation of the quantity of water-vapor present in the air is an important problem. The quantity present depends upon a variety of causes, the temperature and the direction of the wind being the chief ones. A good chemical method for estimating the water consists iu draw- ing a known volume of air over calcium chloride in a weighed tube. This substance has the power to take up water, as we have repeatedly seen. If the tube is weighed after a certain volume of air has been drawn through it, the increase in weight will show the weight of water con- tained in that volume of air. The quantity of water-vapor present in the air varies between comparatively wide limits. At any given tem- perature the air cannot hold more than a certain quantity. When it contains this quantity it is said to be saturated. If cooled down below this temperature the vapor partly condenses, and appears as water. When a vessel contain- ing ice is placed in the air, that which immediately sur- rounds the vessel is cooled down below the point at which the quantity of water-vapor present would saturate the air, and water condenses on the outside of the vessel. Every one has noticed that on a warm cloudy day more water condenses on such a vessel than on a clear cool day. The water-vapor present in the air has an important effect on man. The inhabitants of countries with moist climates apparently have characteristics which are not generally met with in those who inhabit countries with dry climates. The difference between the effect of moist and that of dry air on an individual is well known. When air which is charged with water-vapor comes in 258 COLLEGE CHEMISTRY. contact with cooler air, the vapor condenses and falls as rain. The method employed for the purpose of estimating the quantity of carbon dioxide in the air consists in drawing a known volume of air over something that has the power to absorb the carbon dioxide, and then determining the increase in weight of the absorbing substance. Potassium or sodium hydroxide is well adapted to this. An apparatus has been constructed in which barium hydroxide, Ba(OH) 2 , is used as the absorbent. When carbon dioxide is passed through a solution of this substance insoluble barium car- bonate, BaC0 3 , is thrown down according to the equation Ba(OH) 2 + C0 2 = BaCO, + H 2 0. This may be filtered off and weighed, and the quantity of carbon dioxide estimated from the results; or, if a known quantity of the hydroxide is taken, the quantity left unacted upon after the experiment can be determined by neutralizing with an acid, the neutralizing power of which has previously been determined with care. The quantity of carbon dioxide present in the air is relatively very small, being about 3 parts in 10,000. It varies slightly according to the locality and season, being greater in cities and in summer than in the country and in winter; and greater in warm countries than in cold. It is as essential to the life of plants as oxygen is to the life of animals. It is not an easy matter to determine the quantities of the other constituents of the air, as the ammonia, organic substances, etc., though there is no difficulty in determin- ing that they are present in very small quantities. Air and Life. — The relations of the air to the most im- portant chemical changes which are taking place upon the earth form one of the most interesting subjects of study. We have had a slight glimpse of the action of oxygen, and of that of carbon dioxide; both are essential to the life. of plants and animals. So, too, the water-vapor acts chemic- ally upon plants, and probably to some extent in the respiration of animals. As regards the nitrogen, this PURE AND IMPURE MR. 259 element is frequently referred to as inert, and as serving the purpose of diluting the oxygen. Inert it undoubtedly is, and there is also no doubt that it dilutes the oxygen, but these statements give a very inadequate conception of the important part played by it in the processes of nature. Nitrogen in some form of combination is an essential con- stituent of plants and animals. The animals get their nitrogenous compounds from the plants, and the plants get theirs partly at least from the soil. By the growth of plants, therefore, nitrogenous compounds are constantly being withdrawn from the soil. When plants and animals undergo decomposition in the soil, the nitrogen contained in them is gradually converted into salts of nitric acid or nitrates, and if the decomposition takes place in the air the nitrogen is converted principally into ammonia. Both in the form of nitrates and of ammonia the nitrogen can be utilized by plants, so that if the plants and animals which have received their nourishment from a certain tract of land should be allowed to decay upon this land after death, and the products thus formed should be uniformly distributed in the soil, the latter would not become exhausted. But the products of the soil are removed, and, therefore, the nitrogen required for the growth of other plants is removed, and the soil becomes unproduc- tive. In order that it may be rendered fertile again, the lost nitrogen must be supplied. , It appears from recent very elaborate experiments that the plants have the power to take up from the air a part of the nitrogen which they need. It has been shown that in this absorption of nitrogen from the air certain minute organisms that exist in the soil play a part. Pure and Impure Air. — Pure air may be defined as air that consists of nitrogen, oxygen, and carbon dioxide in the proportions stated above, together with some water- vapor, ammonia, hydrogen dioxide, and argon, and noth- ing of an injurious nature. It is evident from what has been said that there is constant danger of contamination from natural causes. The most common cause of con- 260 COLLEGE CHEMISTRY. tamination is the breathing of human beings in rooms which are inadequately supplied with air. The breathing process involves the using up of oxygen and the giving off of carbon dioxide and small quantities of organic matter which is undergoing decomposition. If the quantity of oxygen is reduced below a certain limit the air becomes unfit for breathing purposes, and evil effects follow. An ordinary inhalation will not then be sufficient to supply the blood with the oxygen necessary to purify it, and the system will begin to suffer. Headache, drowsiness, and a general sense of discomfort follow. The ill effects of breathing the air of a badly-ventilated room occupied by a number of human beings appear to be due for the most part to the presence of the small quantities of decomposing organic matters which are given off from the lungs with the carbon dioxide and other gases. These act as poisons. They have been thrown off from the lungs because they are unfit for use, and when they are taken back again the normal processes of the body are interfered with. The subject of ventilation has been so thoroughly discussed of late years that great improvement has been made in the arrangements for supplying pure air to dwelling apart- ments and audience halls, but there is still room for improvement. Fortunately, in most buildings there is one source of supply of pure air which is independent of architects' plans. This is the diffusion of gases through the porous materials of which the buildings are con- structed. This diffusion was referred to under the head of Hydrogen (see p. 58). There is also a good deal of ventilation through the cracks and other apertures which are always to be found in our buildings. Under some conditions not thoroughly understood the air becomes contaminated by the decomposition of animal and vegetable matter. Air thus contaminated may cause specific diseases, and some of these are spoken of as being caused by malaria, a word which signifies simply bad air. It has been shown that there are in the air microscopic germs which have the power to develop in the body, and ARGON. 261 then to cause the symptoms which are referred to malaria. Disease germs and germs of other kinds are present in the air in great variety, and they play an important part in connection with the life and health of mankind. Liquid Air. — Attention has heen called to the fact that both oxygen and nitrogen can be liquefied. As air is a mixture of these two gases it follows that it can be liquefied. In brief the method employed consists in subjecting the air to a pressure of two thousand to twenty-five hundred pounds to the square inch and cooling this compressed air down to the ordinary temperature by water. It is then passed through a tube that ends in a needle-valve. It is allowed to escape through this valve into another tube that surrounds the one from which it has just escaped. In expanding it is cooled down below the temperature at which the compressed air in the inner tube becomes liquid. This liquid escapes with the gaseous air and is easily col- lected. Liquid air is a turbid, colorless liquid. The turbidity is due to the presence of solid water and solid carbon dioxide. By passing the liquid through a paper filter the solids are removed, and a transparent liquid is thus obtained. This consists mostly of nitrogen and oxygen in the proportion of about four-fifths of the former to one- fifth of the latter. This mixture boils at — 191°. As the nitrogen boils at a lower temperature (— 194°) than oxygen (— 181°), more nitrogen than oxygen is converted into gas in a given time, and after a time the liquid that is left is much richer in oxygen than ordinary air. Argon. — Lord Rayleigh has shown that a litre of nitrogen prepared from the air by abstracting the oxygen and the small quantities of other substances known to be present weighs 1.2572 grams, while a litre of nitrogen made from some chemical compound, such as ammonia, weighs 1.2505 grams. Chemically-prepared nitrogen is lighter than that obtained from the air. This observation led Lord Rayleigh and W. Ramsay to a more thorough chemical examination of the air, the result of which was the discovery of a con- 262 COLLEGE CHEMISTRY. stituent previously unknown. This is the gas argon. It can be obtained by either of two methods: (1) By passing air over heated copper until all the oxygen is abstracted, and then over heated magnesium which unites with the nitrogen but leaves the argon; (2) by mixing air with oxygen and passing electric sparks through the mixture. The oxygen and nitrogen combine and the product can easily be removed. After all the nitrogen has thus been removed argon remains behind. Argon is an element with the atomic weight about 40. It is present in the air to the extent of about 1 per cent of the nitrogen. It cannot be made to combine with any other element. Other Gases in the Air. — By allowing the greater part of liquid air to evaporate and examining the residue Ramsay has discovered another gaseous substance in the air which he calls krypton. This has since been obtained in larger, and probably purer, condition by Ladenburg and Kruegel. It has the specific gravity 58.67, or about 59 (H = 1). By liquefying argon by means of liquid air, and collect- ing the gas given off from, it at successive stages, Ramsay has shown, further, that two other gases, which he calls neon and xenon, are present in the air in very small quantities. EXPERIMENTS. Preparation' of Nitrogen'. Experiment 115. — Place a good-sized stoppered bell-jar over water in a pneumatic trough. In the middle of a flat cork about three inches in diameter fasten a small porcelain crucible, and place this on the water in the trough. Put in it a piece of phos- phorus about twice the size of a pea, and set fire to it. Quickly place the bell-jar over it. At first some air will be driven out of the jar. The burning will continue for a short time, and then gradually grow less and less active, finally stopping. On cooling, it will be found that the volume of gas is less than four-fifths the original volume, for the reason that some of the air was driven out of the vessel at the beginning of the experiment. Before re- EXPERIMENTS: ANALYSIS OF AIR. 263 moving the stopper of the bell-jar see that the level of the liquid outside is the same as that iuside. Try the effect of introducing successively several burning bodies into the nitrogen, — as, for example, a candle, a piece of sulphur, phosphorus, etc. Experiment 116. — Pass air slowly over copper contained in a tube heated to redness and collect the gas which passes through. Does it act like nitrogen ? \ Experiment 117. — In a good- sized Woulff's bottle provided with a safety-funnel and de- livery-tube as shown in Fig. 48, put some copper-turnings and pour upon them concen- trated ammonia, but not enough to cover them. Close the delivery-tube by means of a pinch-cock ; and let the ves- sel stand, What evidence of action is there ? After a time, force some of the gas out of the bottle by pouring water through the funnel, and opening the delivery-tube. Does the gas act like nitrogen ? % Analysis of Air. Experiment 1 18. — Arrange an apparatus as in Fig. 11. In- stead of a plain tube, use one graduated in cubic centimetres. Enclose 60 to 80 cc. air in the tube over water. Arrange the tube so that the level of the water inside and outside is the same. Note the temperature of the air and the height of the barometer. Reduce the observed volume to standard conditions. Now intro- duce a piece of phosphorus, as in Experiment 26, and allow it to stand for twenty-four hours. Draw out the phosphorus. Again arrange the tube so that the level of the water inside is the same as that outside. Make the necessary corrections for temperature, pressure, and the pressure of aqueous vapor. It will be found that the volume has diminished considerably, but that about four fifths of the gas originally put in the tube is still there. If the Fio. 48. 264 COLLEGE CHEMISTRY. work is done properly, the volume of the gas left in the tube will be to the total volume used as 79 to 100. In other words, of every 100 cc. air used 21 cc. are absorbed by phosphorus, and 79 cc. are not. The gas absorbed is oxygen, identical with the oxygen made from the oxide of mercury, manganese dioxide, and potassium chlorate. The gas left over has no chemical prop- erties in common with oxygen. Carefully take the tube out of the vessel of water, closing its mouth with the thumb or some suit- able object to prevent the contents from escaping. Turn it with the mouth upward, and introduce into it a burning stick. Does it support combustion ? Is it oxygen 1 Experiment 119. — Expose a few pieces of calcium chloride on a watch-glass to the air. It gradually becomes liquid by absorb- ing water from the air. Experiment 120. — Expose some clear lime-water to the air. It soon hecomes covered with a white crust. A similar change takes place if baryta- water is exposed in the same way. Lime- water is made by putting a few pieces of quicklime in a bottle and pouring water upon it. The mixture is well shaken up and allowed to stand. The undissolved substance settles to the bot- tom, and with care a clear liquid can be poured off the top. This is lime-water, which is a solution of calcium hydroxide, Ca(OH) 9 , in water. Baryta-water is a solution of a similar com- pound of the element barium. When these solutions are exposed to nitrogen or oxygen, or to an artificially prepared mixture of the two gases, no change takes place. Further, if air is first passed through a solution of caustic soda it no longer has the power to cause the formation of a crust on lime-water or baryta- water. a- Experiment 121. — Arrange an apparatus as shown in Fig. 49. The wash- cylinders A and B are half filled with a solution of ordi- nary caustic soda. The bottle O is filled with water. The tube B, which should be filled with water and provided with a pinch- cock, acts as a siphon. Open the pinch-cock and let the water flow slowly out of the bottle. As it flows out air will be drawn in through the caustic soda in the wash-cylinders. When the bottle is a quarter filled with air pour some water in again until it is full. Then draw all the water off. Now remove the stopper from the bottle, pour in 20 to 30 cc. lime-water and cork the- bottle. The crust formed on the lime-water will now be hardly, if at all, perceptible. There is, therefore, something present in the air under ordinary circumstances which has the power to form a crust on lime-water or baryta-water, and which can be EXPERIMENTS : ANALYSIS OF AIR. 265 removed by passing the air through caustic soda. Thorough examination has shown that this is the compound which chemists call carbon dioxide, and which is commonly known as carbonic acid gas. It is the substance which was obtained by burning charcoal in oxygen. Fio. 50. Experiment 122. — Into the bottle containing the air from which the carbon dioxide has been removed hold a burning stick 266 COLLEGE CHEMISTRY. or taper for a moment. Notice whether a crust is. now formed on the lime-water. Wood and the material from which the taper is made contain carbon. Explain the formation of the crust on the lime-water after the stick of wood or taper has burned for a short time in the vessel. Experiment 123. — Arrange an apparatus as shown in Eig. 50. The bottle A contains air ; B contains concentrated sulphuric acid ; contains granulated calcium chloride ; D is carefully dried and contains a few pieces of granulated calcium chloride and air. Pour water through the funnel-tube into A, when the air will be forced through B and C and into D. But in passing through B and C the moisture contained in it will" be removed, and the air which enters, D will be dry. After A has once been filled with water, empty it and fill it again, letting the dried air pass into D. This operation may be repeated auy number of times. The calcium chloride in D will not grow moist. CHAPTER XVI. COMPOUNDS OF NITROGEN WITH HYDROGEN.— WITH HYDROGEN AND OXYGEN.— WITH OXYGEN ETC. General Conditions which give Rise to the Formation of the Simpler Compounds of Nitrogen We have seen that at ordinary temperatures nitrogen is an inactive element, showing little tendency to combine with other elements. It is nevertheless an easy matter to get com- pounds of nitrogen with many other elements, and among these compounds, some of those which it forms with hydrogen and oxygen are of high importance. When a compound that contains carbon, hydrogen, and nitrogen, and is not volatile, is heated in a closed vessel, so that the air does not have access to it, the nitrogen passes out of the compound, not as nitrogen, but partly in com- bination with hydrogen, in the form of the compound ammonia. Nearly all animal substances contain carbon, hydrogen, oxygen, and nitrogen in many forms of com- bination, some of which are quite complicated. Many of these give off ammonia when heated. Similarly, com- pounds containing carbon, oxygen, and hydrogen, even though they may be thoroughly dry, when heated give off oxygen in combination with hydrogen in the form of water. Both these kinds of decomposition, that which gives ammonia and that which gives water, are to be ascribed to the fact that the compounds of carbon which are heated are unstable at higher temperatures, and when they are broken down the elements contained in them arrange themselves in combination in stable forms, such as the comparatively sinmle compounds, water and am- 267 268 COLLEGE CHEMISTRY. monia. An illustration of this kind of action was referred to in speaking of the preparation of nitrogen by heating ammonium nitrite. This compound breaks down very easily under the influence of heat, the hydrogen and oxygen combining to form the stable compound water, while the nitrogen remains uncombined. Some animal substances, as, for example, urine, give off ammonia when they undergo spontaneous decomposition in the air. This decomposition is generally, if not always, due to the action of minute organisms, the germs of which are in the air, which develop when they come in contact with certain substances. The coal which is used for making illuminat- ing gas contains some hydrogen and nitrogen in chemical combination, and when the coal is heated ammonia is given off with the other products. When animal substances undergo decomposition in the presence of basic compounds where the temperature is . comparatively high, the nitrogen combines with oxygen and with the metal of the base. Either a salt of nitric acid, HNO s , or of nitrous acid, HN0 2 , is formed. In some countries where the conditions are favorable to the process, immense quantities of nitrates are found, chiefly potassium nitrate, KN0 3 , and sodium nitrate, NaN0 3 . Nitrates are, however, found everywhere in the soil. The change of animal and vegetable nitrogenous substances to the form of nitrates is caused by the action of minute living organisms, which are found everywhere, and serve an im- portant purpose in converting the waste animal and vege- table matter into simple compounds that can be utilized by plants. How they effect the change is not known. From the salts of nitric acid which are found in nature, nitric acid itself can easily be made. Nearly all the compounds of nitrogen with which we have to deal are made either from ammonia or from nitric acid. Relations between the Principal Compounds of Nitro- gen. — In studying the compounds of sulphur, we saw that whenever a compound of sulphur is oxidized with a strong COMPOUNDS OF NITROGEN. 269 oxidizing agent the final product of the action is sulphuric acid; and that, on the other hand, under the influence of strong reducing agents these compounds yield hydrogen sulphide as the final product. So, also, the limit of oxidiz- ing action in the case of chlorine is perchloric acid, and of reducing action, hydrochloric acid. Most of the other compounds of sulphur with hydrogen and oxygen are products intermediate between the two limits, sulphuric acid and hydrogen sulphide, as the other compounds of chlorine with hydrogen and oxygen are intermediate between hydrochloric acid and perchloric acid. The limit of reduction of nitrogen compounds is ammonia, NH 3 , and of oxidation, nitric acid, HN0 3 . As has been noticed, the valence of nitrogen towards oxygen is greater than it is towards hydrogen, but the difference is not as marked as in the case of the members of the chlorine group, and in that of the sulphur group. Its hydrogen valence i3 3, its maximum oxygen valence is 5. A tabular list of the oxides of nitrogen and the acids formed by it is here given: Oxides. Acids. Nitrous oxide, N 2 Hyponitrous acid, N 2 (OH) 2 Nitric oxide, NO Nitrous acid, NO(OH) Nitrogen trioxide, } -^ ^ Nitric acid, N0 2 (OH) (Nitrous anhydride), f 2 » Nitrogen peroxide, N0 2 (N 2 4 ) Nitrogen pentoxide, ) ^ n (Nitric anhydride), ( iN * u 5 Besides the above there are the basic compounds, Jiydroxylamine, NH 2 (OH), ammonia, NH 3 , and hydrazine, N a H 4 ; and with water ammonia probably forms the hy- droxide NH 4 (OH), which is a base. In addition to these there is a compound of nitrogen with hydrogen of the formula N 3 H. This is an acid. (See Triazoic Acid.) There is another acid of the same composition as hypo- nitrous acid, but differing from it in properties, the -rela- tion of which to hyponitrous acid is not yet quite clear. 270 COLLEGE CHEMISTRY. With the members of the chlorine group nitrogen forms a few compounds which are characterized by marked instability. So unstable are they that they explode violently when simply touched. The chloride of nitrogen explodes with great violence when the direct rays of the sun are allowed to shine upon it. With sulphur, nitrogen forms two compounds. With sulphur, hydrogen, and oxygen, however, nitrogen forms a number of compounds, one of which has already been referred to in connection with the manufacture of sulphuric acid. This is the so-called nitrosyl-sulphuric acid which is formed by the action of a mixture of the peroxide, N0 3 , and the monoxide, NO, on sulphurous acid, oxygen, and water. In studying the compounds of nitrogen, it will be best to begin with the end products, ammonia and nitric acid. Ammonia, NH S . — The conditions under which ammonia is formed have been mentioned. The chief source at present is the "ammoniacal liquor" of the gas-works. This is the water through which the gas has been passe4 for the purpose of removing the ammonia, and it contains ammonia, or ammonium hydroxide, NH 4 (OH), in solution. By adding hydrochloric acid to this solution the salt ammonium chloride, NH 4 C1, is formed. This is the well- known substance sal ammoniac. It appears that this name has its origin in the fact that common salt or sodium chloride, NaCl, was formerly called sal armeniacum, and that afterward, through a misunderstanding, ammonium chloride came to be known by the same name which under- went change to the form sal ammoniacum, or sal ammoniac. When the ammoniacal liquor is treated with sulphuric acid, ammonium sulphate is formed. Prom one or the other of these salts it is a simple matter to obtain ammonia. For this purpose it is only necessary to treat the salt with some strongly basic compound, as, for example, potassium or sodium hydroxide, or calcium hydroxide. Thus, when a solution of potassium hydroxide is poured on ammonium chloride or sulphate the strong penetrating odor of am- AMMONIA. 271 monia is at once noticed. The first reaction probably results in the formation of ammonium hydroxide, thus: NH 4 C1 +KOH = NH 4 OH + KOI; (NH 4 ) 2 S0 4 + 2KOH = 2NH 4 OH + K 2 S0 4 ; 2NH 4 C1 + Ca(OH) a = 2NH 4 OH + CaCl 2 ; (NH 4 ) 3 S0 4 + Ca(OH) 2 = 2NH 4 OH + 0aSO 4 . But the ammonium hydroxide breaks down very readily into water and ammonia, which escapes as a gas: NH 4 OH = NH 3 + H a O. In the laboratory ammonia is prepared by treating am- monium chloride with slaked lime or calcium hydroxide. The two are mixed in the proportion of two parts of slaked lime to one of ammonium chloride, placed in a flask and gently heated, when the ammonia is given off at once. It is frequently more convenient to heat a strong aqueous solution of ammonia, such as is found in every chemical laboratory. Such a solution when gently heated readily gives off ammonia. Properties. — Ammonia is a colorless, transparent gas with a very penetrating, characteristic odor. In concen- trated form it causes suffocation. Its specific gravity is 0.59; that is to say, it is but little more than half as heavy as air. A litre of the gas under standard conditions weighs 0.7635 gram. It can easily be compressed to the liquid form by pressure and cold. When the pressure is removed from the liquefied ammonia it passes back to the form of gas, and in so doing it absorbs a great deal of heat. These facts are taken advantage of for the artificial preparation of ice. This application will be clear from the following explanation and Pig. 51. An aqueous solution of ammonia saturated at 0° is brought into the strong iron cylinder A, and then gently warmed, while the vessel B is cooled by cold water. The gas given off from A passes through the bent tubes into B, where it is condensed to a liquid. The cylinder A is now placed in a vessel of cold water, and the water which H 2 COLLEGE CHEMISTRY. Fig. 51. is to be frozen is placed in a cylinder D, and this into the hollow space E in the vessel B. The liquid ammonia passes rapidly into the form of gas ■which is absorbed in the water in A, while at the same time so much heat is absorbed that the water in D is frozen. Ammonia does not burn in the air, but does burn in oxygen with a pale yellow- ish flame. It is absorbed by water in very large quantity. One volume of water at the ordinary temperature dissolves about 600 volumes of ammonia gas, and at 0° about 1000 volumes. The sub- stance with which we commonly have to deal under the name of ammonia is a solution of ammonia in water. It is called "spirits of hartshorn" in common language. The solution has the odor of the gas. It loses all its gas when heated to the boiling temperature. The solution shows a strong alkaline reaction, and has the power to neutralize acids and form salts. The conduct of the solu- tion is, in fact, strikingly like that of sodium and potassium hydroxides, and it is believed that in the solution there are present the ions of a compound of the formula NH 4 (OH), known as ammonium hydroxide, and formed by the direct action of ammonia upon water. If this is true, then the action of ammonia upon acids is to.be explained as fol- lows: Ammonium hydroxide is analogous to potassium hydroxide, but differs from it in that it contains the group of atoms NH 4 in place of the atom K. In some way this group plays in the salts formed by ammonia the same part that the elementary atom potassium plays in the salts of potassium, and just as the latter are called potassium salts, so the former are called ammonium salts. According to this the ammonium salts are salts which contain the group NH 4 , known as the ammonium group, in place of the AMMONIA. 273 .hydrogen of the acids. They are formed by direct com- bination of ammonia with the acids, or by the action of ammonium hydroxide upon acids. The analogy between the action of ammonium hydroxide and that of potassium hydroxide upon acids is clearly shown by the aid of the following equations: K(OH) +HC1 =KC1 + H 2 NH 4 (0H) +HC1 = NH 4 C1 + H 2 K(OH) + HN0 3 = KN0 3 + H 2 NH 4 (OH) + HNO s = NH 4 M) 3 + H 2 2K(OH) + H 2 S0 4 = K 2 S0 4 + 2H 2 0; 2NH 4 (OH) + H 2 S0 4 = (NH 4 ) 2 S0 4 + 2H 2 0. The formation of the ammonium salts by direct action of ammonia, NH, , upon acids is represented in the follow- ing equations : NH 3 +HC1 =NH 4 C1; NH 3 + HN0 3 = NH 4 N0 3 ; 2NH 3 + H 3 S0 4 = (NH 4 ) 2 S0 4 . The strong tendency of ammonia to combine directly with acids is shown by bringing two uncovered vessels, one containing a solution of ammonia, and the other a solution of hydrochloric acid, near each other. A dense cloud will at once be formed, if the solutions are concentrated. This is due to the direct combination of the gases which escape from the solutions. While the assumption of the existence of the group ammonium, NH 4 , in the ammonium salts is of great service in dealing with these salts, and while this assump- tion appears to be entirely justified by the facts, no com- pound of this composition has as yet been isolated. The name ammonium is given to the hypothetical compound on account of the fact that it evidently plays the part of a metallic element, and it is customary to give such elements names ending in turn. While, further, it is generally believed that the ions of the compound ammonium 274 COLLEGE CHEMISTRY. hydroxide, NH 4 (OH), are formed when ammonia dissolves in water, the compound itself has not been isolated, owing to its instability and tendency to break down into ammonia and water. On the other hand, some very interesting derivatives of this hydroxide have been isolated. There is one of these which is derived from the hydroxide by the replacement of the four hydrogen atoms of the ammonium by groups of carbon and hydrogen atoms. This compound is stable and can be isolated as a hydroxide of the general formula NR 4 (OH), in which R 4 represents the groups of carbon and hydrogen atoms. In solution it acts almost "^exactly like potassium hydroxide. 4' Composition of Ammonia. — By oxidation under the proper conditions it is possible to convert the hydrogen of ammonia into water and leave the nitrogen in the free state. As water and nitrogen are the only products formed, and the quantity of oxygen used up in the oxidation is equal to the quantity of oxygen found in the water formed, it follows that nitrogen and hydrogen are the only elements contained in ammonia. When electric sparks are passed for some time through a mixture of nitrogen and hydrogen, some ammonia is formed. Conversely, when electric sparks are passed for a time through ammonia, nitrogen and hydrogen are obtained. If, in the oxidation of a known weight of ammonia, the water formed and the nitrogen left uncombined are accurately determined, it will be found that in ammonia the elements are combined very nearly in the proportion of 14 parts by weight of nitrogen to 3 parts by weight of hydrogen. Further, the molecular weight determined by the method of Avogadro is approximately 17. Therefore, the molecular formula of ammonia is NH 3 , the atomic weight of nitrogen being 14. The proportion by volume in which the two elements combine can be determined by the following method: A glass tube closed at one end and provided with a glass stop-cock at the other is filled with pure chlorine gas. By COMPOSITION OF' AMMONIA. 275 means of a small funnel attached to the open end a little of a strong aqueous solution of ammonia is slowly intro- duced into the tube. Reaction takes place at once between the chlorine and the ammonia, according to the equation NH, + 301 = N+8H01; and the hydrochloric acid unites with ammonia to form ammonium chloride: 3HC1 + 3NH 3 = 3JSTH 4 C1. The entire change is therefore represented by the equation 4NH 3 + 3C1 = N + 3NH 4 C1. Hydrogen and chlorine combine in equal volumes, as we hare already learned. Now, if we start with a measured volume of chlorine, and add ammonia to it until it is all used up, we know that the volume of hydrogen which has been extracted from ammonia is equal to the volume of chlorine with which we started. If we measure the volume of nitrogen left over, we know the volume of the nitrogen which was in combination with a volume of hydrogen equal to that of the chlorine originally taken. This experiment has been tried repeatedly, and it has been found that the ratio of the volume of nitrogen to that of the hydrogen" with which it was combined is as 1 to 3. The tube being full of chlorine at the beginning of the experiment, it will be found to be one-third full of nitrogen at the end. Therefore, in ammonia 1 volume of nitrogen is combined with 3 volumes of hydrogen. The experiment just referred to will perhaps be better understood by the aid of the accompanying diagram. The chlorine in the tube may be represented as made up of three equal parts or volumes. Each volume of chlorine combines with an equal volume of hydrogen, leaving the nitrogen uncombined. The volume of nitrogen left is only 276 COLLEGE CHEMISTRY. one-third of that of the chlorine, or for three volumes of chlorine there is one volume of nitrogen: lvol. CI lvol. CI lvol. CI combine with lvol. H lvol. H lvol. H leaving lvol. N Therefore in ammonia the gases nitrogen and hydrogen are combined in the proportion of 1 volume of nitrogen to 3 volumes of hydrogen : lvol. H lvol. H lvol. N = Volume relations in ammonia. lvol. H Another question in regard to the volume relations remains to be answered, and that is: When nitrogen and hydrogen unite in the proportions above stated, how many volumes of ammonia gas do the four volumes of the con- stituents form ? It is not possible to determine this by direct combination of the two gases, but ammonia can be decomposed into its constituents by continued passage of electric sparks through it. When this is done it is found that after the decomposition the gases occupy twice the volume which was occupied by the ammonia. It appears, AMMONIUM AMALGAM. 277 therefore, that -when hydrogen and nitrogen combine to form ammonia the volume is reduced to one-half, or, what is the same thing, when three volumes of hydrogen com- bine with one volume of nitrogen the four volumes* form two volumes of ammonia gas. The above facts have already been commented upon in speaking of the combination of gases in general; and it has been shown that chlorine, oxygen, and nitrogen com- bine with hydrogen in entirely different ways : (1) 1 volume of chlorine combines with 1 volume of hydrogen to form 2 volumes of hydrochloric acid gas. (2) 1 volume of oxygen combines with 2 volumes of hydrogen to form 2 voltimes of water-vapor. (3) 1 volume of nitrogen combines with 3 volumes of hydrogen to form 2 volumes of ammonia gas. What the cause of these differences is we do not know. In some way these facts are directly connected with the law of Avogadro that equal volumes of all gases contain the same number of molecules, and with the power of the atoms of chlorine, oxygen, and nitrogen to combine with one, two, and three atoms of hydrogen respectively. When one atom of chlorine unites with one atom of hydrogen the result is a molecule. So also when one atom of oxygen unites with two atoms of hydrogen, and when one atom of nitrogen unites with three atoms of hydrogen, the result in each case is a molecule, and, according to the law of Avogadro, a gaseous molecule, whether it consists of one atom or a hundred atoms, occupies the same space. Ammonium Amalgam.— A very curious substance which appears to consist of mercury and ammonium is formed when a solution of ammonium chloride is treated with a compound of sodium and mercury known as sodium amalgam. The action is thought to take place thus : 2NH.C1 + Na 2 Hg = 2Na01 + (NH 4 ) 2 Hg. The product, ammonium amalgam, is unstable, break- ing down very soon into ammonia, hydrogen, and mercury. 278 COLLEGE CHEMISTRY. It will be referred to again and somewhat more fully under the head of Mercury. Hydrazine, N 2 H 4 . — A compound closely related to am- monia* and ammonium has recently been prepared by a complicated method that cannot be explained here. This is known as hydrazine. Its composition and molecular weight are represented by the formula N 2 H 4 . A large number of derivatives of hydrazine are known, and have been studied exhaustively. Hydrazine is a liquid that boils at 113.5° in a current of hydrogen. It acts upon acids much as ammonia does, forming the hydrazine salts. Hydroxylamine, KH 2 (OH). — This compound is pre- pared by reducing nitric acid : N0 2 (OH) + 6H = NH,(OH) + 2H 2 0. It can also be prepared by reduction of nitric oxide : 2NO + 6H = 2NH 2 (OH). It is a solid consisting of leaflets or hard needles. It melts at 33.05°, and boils at 58° under a pressure of 22 mm. When the water solution is evaporated, both ammonia and hydroxylamine pass over with the water. Its salts are easily obtained by treating the solution with acids : NH 2 (OH) + HC1 = NH,(OH)Cl; NH 2 (OH) + HN0 3 = NH 3 (OH)N0 3 . From the method of formation and the composition it appears that these salts are ammonium salts in which hydroxyl is substituted for one hydrogen of the ammonium. They should therefore be called hydroxyl-ammonium salts. One of the most characteristic properties of hydroxylamine is the ease with which it breaks down into ammonia, nitrogen, and water: 3NH 2 OH = N a + NH 3 + 3H 3 0. TRIAZOIC ACID— NITRIC ACID. 279 If brought in contact with compounds capable of reduc- tion, it reduces them, the nitrogen in these cases generally combining with oxygen to form nitrous oxide, N 2 0. The reduction of cupric oxide, CuO, takes place according to the equation 2NH 2 (OH) + 4CuO = N 2 + 2Cu 2 + 3H 2 0. By nascent hydrogen hydroxylamine is reduced to ammonia. Triazoic Acid, N 3 H. — This compound, which is also called hydr azoic and hydronitric acid, can be made by a number of reactions inyolving the use of complex organic substances. A simpler method consists in passing am- monia gas over heated metallic sodium, and then passing nitrous oxide, N 2 0, over the resulting product, which is sodium amide. The following reaction takes place : NH 2 Na + N 3 = N 3 Na + H 2 0. By dissolving in water the sodium salt thus formed, treating with dilute sulphuric acid; and distilling, a solu- tion of the acid is obtained. The compound is a colorless liquid that boils at 37°. It has a very penetrating odor, and produces bad effects upon one who inhales it. It is an acid resembling hydrochloric acid. It, as well as some of its salts, is extremely explosive. Its ammonium salt formed by direct union with ammonia is interesting as it has the composition N, NH 4 or N 4 H 4 . Nitric Acid, HN0 3 . — This important chemical compound was first made, though not in pure condition, about the ninth century by distilling saltpetre, copper sulphate, and alum. The name nitric acid has its origin in the fact that the compound is formed from nitre. It has already been stated that the salts of this acid, particularly the potassium and sodium salts, occur very widely distributed in the earth, and that there is a great accumulation of the sodium salt in South America, whence the name Chili saltpetre. Wherever organic matter, particularly that of animal 280 COLLEGE CHEMISTRY. origin, undergoes spontaneous decomposition in the pres- ence of basic substances, nitrates are formed, in conse- quence of tbe action of an organism known as the nitrifying ferment. This process of nitrification has already been referred to in a general way. It is one of great importance for the welfare of the human race, and indeed of most living beings, as by its aid the useless nitrogenous sub- stances of dead plants and animals are converted into the useful nitrates which in the soil aid the process of plant growth. Nitric acid can be formed by the action of electric sparks on nitrogen and oxygen in the presence of water. It is also formed by the action of oxidizing agents on ammonium compounds. Nitric acid is always prepared by treating potassium or sodium nitrate with concentrated sulphuric acid; When sodium nitrate is treated with sulphuric acid action takes place thus : NaN0 3 + H 2 S0 4 = NaHSO, + HNO a . The salt formed in this way is primary or acid sodium sulphate. If sufficient of the saltpetre is present and the temperature is raised, a second reaction takes place, result- ing in the formation of normal sodium sulphate : NaN0 3 + NaHS0 4 = Na 2 S0 4 + HNO s . But the temperature required for this reaction is so high that a considerable part of the nitric acid is decomposed. In the preparation of nitric acid, therefore, the first reaction is the one used, and for this purpose the sub- stances are brought together in retorts in the proportion of their molecular weights (about equal weights), and the ■retort gently heated, The nitric acid distils over slowly, and is condensed by cooling the receiver. On the large scale the acid is made by bringing Chili saltpetre and concentrated sulphuric acid together in cast- iron cylinders or retorts. NITRIC ACID. 281 Nitric acid is a colorless volatile liquid. It begins to boil at 86°, but at this temperature it undergoes partial decomposition into nitrogen peroxide, water, and oxygen : 2HNO3 = 2N0 2 + H 2 + 0. It undergoes the same change slowly when exposed to the direct rays of the sun. In consequence of this decomposi- tion the distillate collected in the manufacture of nitric acid, and, in general, whenever the acid is distilled, always contains a considerable percentage of water, and is colored more or less yellow by the nitrogen peroxide present. In order to abstract the water from ordinary nitric acid it is mixed with concentrated sulphuric acid and slowly dis- tilled; but even under these circumstances the product is colored in consequence of some decomposition, and it also contains some water. By conducting carbon dioxide gas through the gently warmed acid the nitrogen peroxide can be removed, and in this way an acid containing about 99.5 per cent of the compound HN0 3 has been obtained. Pure nitric acid is a very active substance chemically. It gives up its oxygen readily and is itself thus reduced to other compounds of nitrogen and oxygen, or of nitrogen, oxygen, and hydrogen, as has already been pointed out. When it acts upon the metals it forms nitrates, metal atoms being substituted for the hydrogen. According to the conditions, nitrogen peroxide, N0 2 , nitrous acid, HN0 2 , nitric oxide, NO, nitrous oxide N 2 0, nitrogen, hydroxyl- amine, NH 2 (0H), and, finally, ammonia are formed by reduction of the acid. Of these reactions, that which gives nitric oxide, NO, is the one which commonly takes place on treating metals with nitric acid. The oxides N0 2 and N 2 3 are themselves readily reduced to nitric oxide. If the element upon which the . acid acts has not the power to take the place of the hydrogen, the action con- sists in oxidation. This is shown in the action of strong nitric acid upon tin, phosphorus, carbon, sulphur, etc. In each case the highest oxidation-product is formed. Tin is converted into normal stannic acid, Sn(0H) 4 ; phos- 282 COLLEGE CHEMISTRY. phorus into phosphoric acid, PO(OH) 3 ; carbon into carbon dioxide, C0 2 ; and sulphur into sulphuric acid, S0 2 (OH) 2 . It disintegrates carbon compounds very readily, converting them into their final products of oxidation. In contact with the skin it causes bad wounds. The acid mostly used in the laboratory has the specific gravity 1.2 and contains 32 per cent nitric acid, HNO s . The commercial acid contains about 68 per cent of the acid. When a mixture of nitric acid and water is boiled under the ordinary atmospheric pressure it loses either water or nitric acid until it contains 68 per cent of the acid and then it distils over. This does not correspond to any definite hydrate of nitric acid, though it approximates the composition required by normal nitric acid, N(OH) 5 , or HN0 3 -f- 2H 2 0, and it is probable that this hydrate is the chief constituent of the mixture. Nitric acid is a strong monobasic acid, forming salts of the general formula MN0 3 , all of which are soluble in water. Because nitric acid is a strong acid, and all normal nitrates are soluble in water, it is one of the best solvents. Red Fuming Nitric Acid is formed in the manufacture of nitric acid from saltpetre and sulphuric acid if the tem- perature is raised to a sufficient extent to cause the acid sulphate to act upon the nitrate : NaN0 3 + HNaSO, = Ka 2 S0 4 + HNO s . At this temperature the nitric acid undergoes considerable decomposition. The nitrogen peroxide formed is absorbed by the nitric acid, and the product thus obtained is the red fuming acid. It acts more energetically than nitric acid, and finds some applications in the laboratory and in the arts. When heated it gives off nitrogen peroxide, and if diluted with water it is changed to ordinary nitric acid, as nitrogen peroxide is decomposed by water, forming nitric acid and nitric oxide or nitrous acid, according to the temperature of the water. NITROUS ACID. 283 Nitro-hydrochloric Acid or Aqua Regia is a liquid formed by mixing concentrated nitric and hydrochloric acids. It was called aqua regia because it can dissolve gold, the king of the metals. The active power of this liquid as a solvent of metallic substances is due to the fact that it gives off chlorine, and a compound of nitrogen, oxygen, and chlo- rine which readily gives up its chlorine. This compound has the composition represented by the formula NO CI, and it is best designated by the name nitrosyl chloride. The product of the action of nitro-hydrochloric acid upon a metal is the corresponding chloride. Nitrous Acid, HN0 2 . — When certain salts of nitric acid are reduced they yield the corresponding nitrites. Thus, when potassium nitrate is heated with metallic lead this reaction takes place : KN0 3 + Eb = KN0 2 + PbO. Indeed, if potassium nitrate is heated alone it loses oxygen and is converted into the nitrite : 2KN0 3 = 2KN0 2 + 2 ; but the reaction is not complete, and the salt thus obtained always contains more or less nitrate. Nitrous acid is known only in solution. If an attempt is made to isolate it from a nitrite the product is the anhy- dride, nitrogen trioxide, N 2 3 . Thus, if sulphuric acid is added to potassium nitrite the following reaction takes place : 2KNO, + H 2 S0 4 = N 2 3 + H 2 + K,SO, It may be that the first action is the liberation of nitrous acid, and that this then breaks down by loss of water. The two reactions are represented thus : 2KN0 2 + H 2 SO, = 2HN0 2 + K 2 SO,; 2HN0 2 = N 2 3 + H 2 0. The trioxide, N,0,, however, breaks down into the two gases nitrogen tetroxide, N0 2 , and nitric oxide, NO. 284 COLLEGE CHEMISTRY. A certain analogy will be observed between this action and that which takes place in the action of potassium hydroxide upon an ammonium salt, when ammonium hydroxide is probably first given off, and then breaks down into ammonia and water. Salts are known which are derived from normal nitrous acid, N(OH) s , but most of the nitrites are derived from the acid of the formula NO(OH), which is to be regarded as formed from the normal acid by loss of one molecule of water : N(OH) 3 = NO (OH) + H 2 0. Hyponitrous Acid, H 2 N 2 2 . — The sodium salt of this acid is made by reducing sodium nitrite in solution by means of sodium amalgam : 2NaNO s + 8H = Na 2 N 2 2 + 4H 2 0. The acid can also be made by oxidation of hydroxylamine. It is a solid consisting of white crystalline plates. It is very explosive when freed from water. In water solution, it is much more stable, but at the ordinary temperature it breaks down gradually, the principal products being nitrous oxide and water: H 2 N 2 2 = N 2 + H 2 0. Nitrous Oxide, N 2 0. — -This compound can be obtained by reduction of nitric acid, and is sometimes formed in considerable quantity when copper is treated with the concentrated acid, though when made in this way it is always mixed with a large proportion of nitric oxide. The best way to make it is to heat ammonium nitrate, NH 4 N0 3 , which breaks down into nitrous oxide and water: NH 4 N0 3 = N 2 + 2H 2 0. In the same way we have seen that ammonium nitrite breaks down into free nitrogen and water when heated : NH 4 NO, = N 2 + 2H 2 0. NITROUS OXIDE— NITRIC OXIDE. 285 In these reactions we see exhibited the tendency of hydrogen and oxygen to combine at elevated temperatures. At ordinary temperature this tendency is not strong enough to cause a disturbance of the equilibrium of the parts of the compound. As the temperature is raised and the equilibrium thus disturbed, the affinity of the hydrogen for the oxygen asserts itself. The two elements combine to form water, and the decomposition above represented takes place. Nitrous oxide is a colorless, transparent gas which has a sweetish taste and odor. Its specific gravity is 1.53. It is somewhat soluble in water; one volume of water at 0° dissolving somewhat more than its own volume of the gas. It supports combustion almost as well as pure oxygen. Some substances which burn in oxygen do not, however, burn in nitrous oxide. Sulphur which burns in oxygen is extinguished in nitrous oxide, unless it is previously heated to a high temperature. When inhaled, nitrous oxide causes a kind of intoxica^ tion, which is apt to show itself in the form of hysterical laughing. Hence the gas is called laughing gas. Inhaled in larger quantity it causes unconsciousness and insensi- bility to pain. It is therefore used extensively to prevent pain in some surgical operations, particularly in extracting teeth. "When subjected to a low temperature and high pressure the gas is easily liquefied, and enclosed in properly con- structed metallic cylinders the liquid is now sent into the market. In order to get the gas it is only necessary to open the stop-cock of the cylinder. When the liquid comes in contact with the air it rapidly turns to gas, and the temperature is very much lowered in consequence. This causes a part of ihe liquid to solidify. Nitric Oxide, NO. — This is the most stable compound of nitrogen and oxygen, and is the most common product of the reduction of nitric acid. Thus, when nitric acid acts upon copper and other metallic elements the chief product is generally nitric oxide, though, as we have seen, 286 COLLEGE CHEMISTRY. the reduction may be carried farther. The principal action in the case of copper is represented thus : 8HNO s + 3Cu = 3Cu(N0 3 ) 2 + 2NO + 4H a O. Considering the ease with which nitric acid gives up its oxygen, and the ease with which copper takes up oxygen, it is probable that the copper abstracts oxygen directly from the acid as represented thus : 2KNO3 + 3Cu = CuO + H 2 + 2N0. In this case the copper oxide would at once form copper nitrate with the excess of nitric acid : 6HN0 8 + 3CuO = 3Cu(N0 s ) 2 + 3H 2 0. Or, combining the two equations, the total action is repre- sented in the same way as it is above. The nitric acid must not have a specific gravity higher than 1.2, and the temperature must be kept down, otherwise the reduction of the nitric acid is carried farther and considerable nitrous oxide is formed. It is possible that to some extent the hydrogen liberated from the acid may act as a reducing agent, thus causing the formation of the lower oxides of nitrogen, as, for example, 2HNO, + 6H = 2NO + 4H 2 0. Nitric oxide is a colorless, transparent gas. Its most remarkable property is its power to combine directly with oxygen when the two are brought together. The act of combination is not accompanied by the appearance of light, though heat is evolved. In the reaction which takes place at ordinary temperatures nitrogen peroxide, N0 2 , is formed: NO + = N0 2 . The product is a colored gas, and the change of the color- less nitric oxide to this colored product can therefore easily NITROGEN TRIOXIDE— NITROGEN PEROXIDE. 287 be recognized. This reaction is, further, the chief cause of the reddish-brown fumes seen when nitric acid acts upon metals and other elements. At a low temperature some nitrogen trioxide is formed when oxygen acts upon nitric oxide. From what has already been said, it will appear that in nitric oxide the oxygen and nitrogen are more firmly united than in the other oxides. Most burning substances are extinguished when introduced into it, though a few when heated in it to a high temperature extract all or a part of the oxygen. Zinc and iron extract half the oxygen and convert nitric oxide into nitrous oxide. Potassium and sodium decompose it, leaving the nitrogen free. Nitrogen Trioxide, N 2 3 . — This oxide is formed by addition of oxygen to nitric oxide at low temperatures ; by decomposition of the nitrites by means of acids; and by the combination of nitric oxide with the peroxide at a temperature below — 21°. The gas given off when nitric acid is reduced with starch or arsenious oxide, As 2 3 , appears to be a mixture of nitric oxide and the peroxide. Pure nitrogen trioxide is a liquid of an indigo-blue color. At a temperature below 0° it undergoes partial decomposi- tion into nitrogen peroxide and nitric oxide : N 2 3 = NO + N0 2 . With cold water nitrogen trioxide undergoes decomposition accompanied by an evolution of nitric oxide. Possibly this reaction takes placB : 3N 2 3 + H 2 = 2HN0 3 + 4NO. By treating the oxide with a solution of sodium hydroxide or potassium hydroxide the corresponding nitrite is formed : 2KOH + N 2 3 = 2KN0 2 + H 2 0. Nitrogen Peroxide, N0 2 . — When nitric oxide and oxygen are brought together in the proportion of 2 volumes of the 288 COLLEGE CHEMISTRY. former to 1 volume of the latter they combine completely to form nitrogen peroxide. These relations will be readily understood when it is borne in mind that 2 molecules of nitric oxide require 1 molecule of oxygen to effect the change, as is shown in the equation 2NO + 2 = 2NO 2' The compound is most easily obtained by heating lead nitrate, when nitrogen peroxide and oxygen are given off, and lead oxide remains behind in the vessel: Pb(NO,) 2 = PbO + 2N0 2 + 0. If the gases are passed through a tube surrounded by a freezing mixture the peroxide is condensed to the form of liquid, while the oxygen passes on. When perfectly dry the peroxide is easily solidified. It acts energetically upon compounds that have the power to take up oxygen. When treated with water it undergoes decomposition. If the temperature is low, nitrous and nitric acids are formed: 2N0 2 + H 2 = HN0 2 + HNO s . If the water is hot, however, the products are nitric acid and nitric oxide : 3N0 2 + H 2 = 2HN0 S + NO. The nitric oxide thus formed will take up oxygen from the air and yield nitrogen peroxide again, and this, in contact ■ with hot water, will be decomposed, forming nitric acid and nitric oxide, until all the peroxide is converted into nitric acid. The determinations of the specific gravity of the gas from the peroxide show that at low temperatures the molec- ular formula is N 2 4 , but that when the temperature 150° is reached the molecule is represented by the formula N0 2 . The compound appears therefore to undergo gradual COMPOUNDS OF NITROGEN. 289 decomposition or dissociation by heat, so that until the temperature 150° is reached the gas is a mixture of the compounds N 2 4 and N0 2 . Nitrogen Pentoxide, ]ST 2 5 . — This compound, which bears to nitric acid the relation of an anhydride, is formed by passing chlorine over silver nitrate and condensing the product. The reaction takes place thus: 2AgNO s + 01 2 = N 2 5 + 2AgCl + 0. It is also formed by treating nitric acid with phosphorus pentoxide, P 2 5 , a compound that has a very marked power to unite with water. The action is represented thus : 2HNO3 = N,0, + H 2 0. The pentoxide is a crystallized substance, which readily decomposes into nitrogen peroxide and oxygen. In conse- quence of the ease with which it gives up its oxygen it acts violently upon many oxidizable substances. "With water it forms nitric acid: N 2 5 + H 2 = 2HN0 3 . Compounds of Nitrogen with the Elements of the Chlorine Group. — Notwithstanding the ease with which chlorine combines with most elements, and the stability of the compounds which it forms with them, its compound with nitrogen is extremely unstable. It can be made by the action of chlorine on ammonia, and by decomposing a solution of ammonium chloride by means of an electric current. In the latter case chlorine is liberated at one of the poles and then acts upon the ammonium chloride : NH.01 + 601 = 4HC1 + NC1 3 . It appears that when chlorine acts upon ammonia different products are formed by the substitution of chlorine for the hydrogen atom for atom, thus : 290 COLLEGE CHEMISTRY. NH 3 + Cl 2 = NH 2 C1 + HC1; NH 2 C1 + Cl 2 = NH01 2 + HC1; NH01 2 + Cl 2 = NCI, + HC1. According to this, the trichloride of nitrogen is the final product of the substituting action of chlorine upon am- monia. The compound is an oil, which undergoes decom- position very readily. It is, indeed, one of the most explosive substances known. It is decomposed by heat, and especially by contact with certain substances, among which are oil of turpentine and caoutchouc. It is slowly decomposed by water, though, probably owing to the slight affinity of nitrogen for oxygen, the decomposition does not take place as readily as that of the compounds of sulphur and chlorine. Direct sunlight causes explosion of the chloride. When ammonia is treated with iodine reactions take place similar to those which take place with chlorine. The product is the iodide of nitrogen, N,H 3 I 3 . This com- pound, like the chlorine compounds, is extremely explo- sive. The simplest way to prepare it is to place a little powdered iodine on a filter and pour concentrated ammonia over it. The substance should be made in only very small quantities at a time. "When dried it decomposes with violent explosion by contact even with soft substances. EXPERIMENTS. Preparation and Properties oe Ammonia. Experiment 124. — To a little ammonium chloride on a watch- glass add a few drops of a strong solution of caustic soda, and notice the odor of the gas given off. Do the same thing with caustic potash. Mix small quantities of ammonium chloride and lime in a mortar, and add a few drops of water. " Experiment 125. — Mix 20 parts iron filings, 1 part potassium nitrate, and 1 part solid potassium hydroxide, and heat the mix- ture in a test-tube. Is there any evidence of the formation of ammonia ? ^ Experiment 126.— Arrange an apparatus as shown in Fig. 37, omitting, however, the funnel-tubo ; a cork with one hole will 291 EXPERIMENTS WITH AMMONIA. therefore serve. Weigh 100 grams quicklime in the flask, and add just enough water to slake it without making it moist ; then add 50 grams ammonium chloride, and mix by shaking. Heat on a sand-bath. After the air is driven out, the gas will be completely absorbed by the water in the first Woulff's flask if shaken from time to time. Disconnect the delivery-tube from the series of Woulff's flasks, and connect with another tube bent upward. Collect some of the gas by displacement of air, planing the vessel with the mouth downward. (Why ?) The arrangement is shown in Fig. 52. The vessel in which the gas is collected should be dry, as water absorbs ammonia very readily. Hence, also, it cannot be collected over water. In the gas collected introduce a burning stick or taper. Ammonia does not burn in air, nor does it support combustion. In working with the gas great care must be taken to avoid inhaling it in any quantity. After enough has been collected in cylinders to exhibit the chief properties, connect the delivery-tube again with the series of Woulff's flasks, and pass the gas over the water as long as it is evolved. Ammonia Bursts in Oxygen. Experiment 127. — Put a little of a concentrated solution of ammonia in a flask placed upon a tripod. Heat gently and, from a gasometer, pass a rapid current of oxygen through a bent tube into the liquid. Apply a light to the mouth of the vessel, when the ammonia will be seen to burn. Fig. 52. Ammonia forms Ammonium Salts with Acids. ^ Experiment 128.— Put 100 cc. of a dilute solution of ammonia in an evaporating-dish. Try its effect on red litmus paper. Slowly add dilute hydrochloric acid until the alkaline reaction is destroyed and the solution is neutral. Evaporate to dryness on a water- bath. Compare the substance thus obtained with sal-ammoniac, or ammonium chloride. Taste. Heat on a piece of platinum- foil. Treat with a caustic alkali. Treat with a little concentrated sulphuric acid in dry test-tubes. Do they appear to be identical ? Similarly sulphuric acid and ammonia yield ammonium sulphate; nitric acid and ammonia yield ammonium nitrate ; etc. 2 9 2 COLLEGE CHEMISTRY. Experiment 129. — Fill a dry cylinder with ammonia gas, and another of the same size with hydrochloric acid gas. Bring them together with their mouths covered. Quickly remove the covers, when a dense white cloud will appear in and about the cylinders. This will soon settle on the walls of the vessels as a light white solid. It is ammonium chloride. Thus, from two colorless gases we get a solid substance by an act of chemical combination. Heat is evolved in the act of combination. Composition of Ammonia. Experiment 130.— This experiment should be performed by a person experienced in the use of chemical apparatus. A glass tube, such as represented in Fig. 53, provided with a glass stop- cock is needed. Fill this tube with chlorine free from air over a saturated solution of sodium chloride. After it is filled let it stand for some time mouth downward in the solution of sodium chloride to let the liquid drip out of it. Close the stop-cock and remove the tube from the solution. Hold it mouth up- ward, and pour a concen- trated solution of ammonia into the funnel-like projec- tion above the stop-cock; put in the glass stopper, and now by slightly opening the stop- cock let the ammonia pass drop by drop into the tube. Eeaction between the chlo- rine and the ammonia takes place, accompanied by a marked evolution of heat, and in a partly-darkened room light is seen. Great care must be taken not to admit air with the ammonia. After nearly all the ammonia V.as passed in from the fun- nel, pour into the funnel about two tninte as much ammonia as EXPERIMENTS WITH NITRIC ACID. 2 93 has already been used, and let this in gradually. Leave the stop- cock closed, and fill the funnel with dilute sulphuric acid. Fit a bent tube into a cork; fill this tube with dilute sulphuric acid ; put the cork in the funnel, and the other end of the tube in a small beaker containing dilute sulphuric acid, and, after immers- ing the long tube in water of the ordinary temperature, open the stop-cock. If the operation has been carried out as it should be, the dilute acid will flow into the tube until it is two-thirds full, and will then stop. The residual gas is nitrogen. What evidence in regard to the composition of ammonia is furnished by this experiment ? The arrangement of the apparatus in the last stage of the ex- periment is shown in Fig. 53. Preparation and Properties of Nitric Acid. Experiment 131. — Arrange an apparatus as shown in Fig. 54. In the retort put 20 grams sodium nitrate (Chili saltpetre) and 20 Fig. 54. grams concentrated sulphuric acid. On gently heating, nitric acid will distil over, and be condensed in the receiver. — What re- action takes place ? — After the acid is all distilled off, remove the contents of the retort. Kecrystallize the substance from water, and compare it with the sodium sulphate obtained in the prepara- tion of hydrochloric acid. (See Experiment 74.) In the latter stage of the operation the vessels become filled with a reddish-brown gas. The acid which is collected has a somewhat yellowish color. "V Experiment 132. — Mix together 4o0 grams concentrated sul- phuric acid and 85 grams ordinary concentrated nitric acid. 294 COLLEGE CHEMISTRY. Pour the sulphuric acid into the nitric acid. Distil the mixture from a retort arranged as in the preceding experiment, taking care to keep the neck of the retort cool by placing filter-paper moistened with cold water on it. Use the acid thus obtained for the purpose of studying the properties of pure nitric acid. Nitric Acid gives up Oxygen readily, and is hence a good Oxidizing Agent. -^Experiment 133. — Pour concentrated nitric acid into a wide test-tube, so that it is about one-fourth filled. Heat the end of a stick of charcoal of proper size, and, holding the other end with a forceps, introduce the heated end into the acid. It will continue to burn with a bright light, even though it is placed below the Fig. 55. surface of the liquid. The action is oxidation. The charcoal in this case finds the oxygen in the acid and not in the air. Great care must be taken in performing this experiment. The charcoal should not come in contact with the sides of the test-tube. A large beaker-glass should be placed beneath the test-tube, so that in case it breaks the acid will be caught and prevented from doing harm. The arrangement of the apparatus is shown in Fig. 55. The gases given off from the tube are offensive and poisonous. Hence this experiment as well as all others with nitric acid should be carried on under a hood in which the draught is good. *■ Experiment 134. — Bnil a little strong nitric acid in a test-tube in the upper part of which some horse-hair (or woollen yarn) has been introduced in the form of a stopper. The horse hair (or yarn) will take fire and burn, and leave a white residue. Hold the test- EXPERIMENTS WITH NITRIC ACID. 295 tube with a forceps over a vessel to catch the contents should the x tube break. ^ Experiment 135.— In a small flask put a few pieces of granu- lated tin. Pour on this just enough strong nitric acid to cover it. Heat gently over a small flame. Soon action will take place. Col- ored gases will be evolved, the tin will disappear, and in its place will be found a white powder. This consists mostly of tin and oxygen. (See Experiment 13. Metals dissolve in Nitric Acid, forming Nitrates. "^ Experiment 136. — Dissolve a few pieces of copper-foil in or- dinary commercial nitric acid diluted- with about half its volume of water. The operation should be carried on in a good-sized flask and under an efficient hood. When the copper has disap- peared, pour the blue solution into an evaporating- dish, and evaporate down to crystallization. Compare the substance thus obtained with copper nitrate. Heat specimens of each. Treat small specimens with sulphuric acid. What evidence have you that the two substances are identical ? Nitrates are decomposed bt Heat. - Experiment 137. — Heat some potassium nitrate in a test-tube. Introduce a piece of wood with a spark on it. Heat also lead nitrate, copper nitrate, and any other nitrates that may be available. What difference do you observe between the decom- position of potassium nitrate and that of lead nitrate ? Nitrates are Soluble in Water. Experiment 138. — Try the solubility in water of the nitrates used in the last experiment. Nitric Acid is reduced to Ammonia by Nascent Hydrogen. * • Experiment 139. — In a good-sized test-tube treat a few pieces of granulated zinc with dilute sulphuric acid. What jgejolved ? Prove it. Now add drop by drop dilute nitric acid? The hydro- gen ceases to be given off. Pour the contents of the tube into an evaporatiug-dish and evaporate the liquid. Put the residue into a test-tube and add a solution of caustic soda, when the smell of ammonia will be noticed. Try the action of the gas on red litmus paper. Moisten the end of a glass rod with a little hydrochloric acid, and hold it in the tube. White fumes are seen. What are they? Do the same with nitric acid. What are the fumes in this case ? 296 COLLEGE CHEMISTRY. Nitrous Acid. Experiment 140. —Melt 25 grams potassium nitrate in a shal- low iron plate and gradually add 50 grams metallic lead cut in small pieces. Stir them together as thoroughly as possible. After the mass is cooled down, break it up and treat with water in a flask. The potassium nitrate will dissolve, while the lead oxide and unused lead will not dissolve. Filter. Add a little, sulphuric acid to some of the solution. A colored gas will be given off. See whether a solution of potassium nitrate acts in the same way. Treat with sulphuric acid a little of the residue left after heating potassium nitrate alone in a test-tube as in Experiment 138. Nitrous Oxide. 1 Experiment 141.— In a retort heat 10 to. 15 grams crystallized ammonium nitrate until it has the appearance of boiling. Do not heat higher than is necessary to secure a regular evolution of gas. Connect a wide rubber tube directly with the neck of the retort and collect the evolved gas over water, as in the case of oxygen. It supports combustion almost as well as pure oxygen. Try experiments with wood, a candle, and a piece of phosphorus. Nitric Oxide. Experiment 142. — Arrange an apparatus as shown in Fig. 56. In the flask put a few. pieces of copper-foil. Cover this with water. Now add slowly, waiting each time for the action to begin, ordinary concentrated nitric acid. When enough nitric acid has been added gas will be evolved. If the acid is added rapidly, it not infrequently happens that the evolution of gas takes place too rapidly, so that the liquid is forced out of the flask through the funnel- tube. This can be avoided by not being in a hurry. At first the vessel becomes filled with a reddish-brown gas, but soon the gas evolved becomes colorless. Collect over water two or three vessels full. The gas collected is prin- cipally nitric oxide, NO, though it is fre- quently mixed with a considerable quantity of nitrous oxide. Experiment 143.— Turn one of the vessels containing colorless nitric oxide with the Fig. 56. mouth upward, and uncover it. The colored gas is at once seen, EXPERIMENTS WITH NITRIC OXIDE, ETC. 297 presenting a very striking appearance. Do not inhale the gas. Perform the experiments with nitric oxide where there is a good draught. Experiment 144. — Pass nitric oxide into a concentrated solu- tion of ferrous sulphate. Afterwards heat the solution and col- lect the gas. What do you conclude that the gas is ? Nitrogen Trioxide. Experiment 145. — In a flask fitted with a safety-funnel and a delivery-tube pour nitric acid of specific gravity 1.30-1.35 upon coarsely-granulated arsenious oxide, As 2 O s . Heat gently, and conduct the gases through a tube surrounded by a freezing mix- ture, as in Experiment 110. Nitrogen Peroxide. Experiment 146. — Admit a little air to nitric oxide contained in a bell-jar over water, and let the vessel stand. Almost imme- diately the color will disappear, showing that the nitrogen perox- ide formed is decomposed. Again admit air, and let the vessel stand. The same changes will be noticed as in the first instance. If oxygen is used instead of air the above changes can be repeated over and over again. Devise an experiment for the purpose of determining whether the nitric oxide is gradually used up or not. CHAPTER XVII. ELEMENTS OF FAMILY V, GROUP B: PHOSPHORUS.— ARSENIC— ANTIMONY.— BISMUTH. THE ELEMENTS AND THEIR COMPOUNDS WITH HYDROGEN. General. — The elements of this group bear to nitrogen much the same relations that the members of the sulphur group bear to oxygen, and those of the chlorine group bear to fluorine. In general they form compounds of the same character and of similar composition. At the same time gradations in properties are noticed in passing from one end of the group to the other. Like nitrogen, the elements of the group are strongly marked acid-formers, though this character grows less marked from nitrogen to bismuth^y Antimony is both an acid -forming and a base- "forming element, while bismuth is more basic than acid. The stability of the hydrogen compounds decreases from nitrogen to antimony; while bismuth does not form a compound with hydrogen. Ammonia, as we have seen, is strongly basic; the corresponding compound of phosphorus and hydrogen has weak basic properties, while those of arsenic and antimony have no basic properties. These hydrogen compounds correspond in composition to am- monia. They are: NH S : PH S AsH 3 SbH s With chlorine they all form compounds corresponding to nitrogen trichloride, and phosphorus and antimony form compounds in which they are quinquivalent, while bismuth 298 ELEMENTS OF FAMILY V, GROUP B. 299 forms a chloride, Bi 2 Cl 4 , in addition to the trichloride. The compounds referred to are : Bi 2 Cl 4 NC1 3 : PCI, AsCl 3 SbCl, Bi01 3 PC1 S SbCl s They all form two oxides corresponding to nitrogen tri- oxide and pentoxide : *fi t : PA As 2 3 Sb 2 3 Bi 2 3 N.O.: p 2 o 5 As 2 6 Sb 2 5 Bi 2 O s None of the elements of the group forms as great a variety of compounds with oxygen as nitrogen does. Antimony, however, forms the oxide Sb 2 4 , corresponding to nitrogen peroxide, N 2 4 ; and bismuth forms the oxide Bi 2 O a or BiO, corresponding to nitric oxide, NO. The hydroxyl compounds or acids, like those of nitro- gen, are related to the maximum hydroxyl compounds of the elements with the valence 5, and to the maximum hydroxyl compounds of the elements with the valence 3. That is to say, they may be regarded as derived from a hydroxide of the general formula M(OH) 6 , and another of the formula M(OH) 3 . Where M is nitrogen these acids break down to the forms N0 2 (OH) and NO(OH) by loss of one or two molecules of water. In the case of the ele- ments of the phosphorus group, however, the breaking down is not generally carried as far as with nitrogen. The general rule is the same as in the sulphur and chlorine groups : the normal acid breaks down to form compounds containing the same number of hydrogen atoms as the hydrogen compounds of the elements. Thus the hydroxyl derivatives of chlorine generally break down to form com- pounds containing one atom of hydrogen, or the same number that is contained in the hydrogen compound, hy- drochloric acid, thus: Cl(OH), yields C10 3 (OH); Cl(OH) 6 yields C10 2 (OH), etc. So, also, in the sulphur group, S(0H) 6 yields S0 2 (OH) 2 , etc., the number of hydrogen 300 COLLEGE CHEMISTRY. atoms in the common form of the acid being the same as that in the hydrogen compound of sulphur, SH 2 . Of the elements of this group phosphorus occurs most abundantly in nature, arsenic and antimony next, and bismuth least abundantly. Arsenic, antimony, and bis- muth occur to some extent in the uneombined condition. Phosphorus occurs in combination. All the elements of the group find applications in the arts, either as the ele- ments or in the form of compounds. Phosphokus, P (At. Wt. 31). Occurrence. — The name phosphorus is derived from the Greek + H 2 0; Normal phosphoric acid Orthopliosphoric acid 320 COMPOUNDS OF THE PHOSPHORUS GROUP. 321 As(OH) 5 Normal arsenic acid - As |(OH) 3 Orthoarsenic acid Sb |(0H) 8 + + H 2 0; H.O. Sb(OH) 5 = Normal antimonic acid Orthoantimonic"acid BismutHc acid appears, however, to be formed from the normal acid by loss of two molecules of water, just as the so-called metaphosphoric, metarsenic, and metantimonic acids are: Bi(OH) 5 Normal bismuthic acid " Bismuthic acid + 2H a 0; P(OH) 5 = r ]OH Metaphosphoric acid + 2H a O; As(OH) 5 = As iSk Metarsenic acid + 2H a 0; Sb(OH) 6 = Sb | OR Metantimonic acid + 2H a 0. Prom the ordinary or ortho acids and from the meta acids more complex forms can be derived by loss of differ- ent quantities of water. The most common form besides those mentioned is that seen in the so-called pyro acids, of which pyrophosphoric acid is the best known example. It is formed from the ortho acid by loss of one molecule of water from two molecules of the acid, just as pyrosulphuric or disulphuric acid is formed from two molecules of ordi- nary sulphuric acid by loss of one molecule of water. The formation of pyrophosphoric acid from orthophosphoric acid takes place according to the equation OH OH OH OH OH OH OH OH OH OH H,0; 3 22 COLLEGE CHEMISTRY. or 2H 3 P0 4 = H 4 P 2 0, + H 3 0. Orthophosphoric acid Pyrophosphone acid Pyroarsenic and pyroantinionic acids bear the same relations to the ortho acids that pyrophosphoric acid bears to orthophosphoric acid. By partial oxidation of phosphorus in presence of water, phosphorous acid, H 3 PO s , is formed. The same acid is formed by the action of phosphorus trichloride on water. According to the latter method of formation we should expect to find that this acid is normal phosphorous acid, P(OH) 3 . It appears probable, however, that the acid has / H the constitution = P^OH. The acids of arsenic and x OH antimony of similar composition seem to be the norma] acids As(OH) 3 and Sb(OH) 3 . The hydroxyl derivative of bismuth corresponding to these acids has no acid proper- ties, but on the contrary is basic. Hypophosphorous acid has the composition H 3 P0 2 . It is monobasic, and it appears therefore that it contains but one hydroxyl, as represented in the formula H 2 OP(OH). It is probable that the relations between phosphoric, phosphorous, and hypophosphorous acids should be represented by the formulas (OH f H ( H OP \ OH, OP \ OH, OP^H . (.OH (oh (OH Phosphoric acid Phosphorous acid Hypophosphorous acid The fundamental compound, then, from which these may be regarded as derived is the unknown oxyphosphine OPH 3 . By oxidation we should expect phosphine to yield in successive stages the three products above named : (H (H (H (H (OH P^H, OP^H, OP^H , OP^OH, OP^OH. (H (H (OH (OH (OH Unknown Hypophosphorous Phosphorous Phosphoric acid acid acid PHOSPHORIC ACID. 323 The oxidation of hydrogen sulphide takes place similarly, as has been shown: s| H , o 2 sj H , o.sJq^ o 2 sj OH . Unknown Sulphurous acid Sulphuric acid With oxygen and chlorine the elements of the phos- phorus group form a number of compounds known as oxychlorides. Towards chlorine as well as towards oxygen all these elements except bismuth are quinquivalent.. A part or all of the oxygen of the oxygen compounds can be replaced by chlorine. Starting with the chlorine com- pound on the one hand, oxychlorides can be obtained from it, until oxygen is substituted for all the chlorine, and the limit is reached in the oxide. So, also, hydroxyl can be substituted for the chlorine and the acids thus obtained. With sulphur phosphorus apparently forms a large number of compounds. Among them are two which have the formulas P 2 S 3 and P 2 S 5 . These plainly are analogous to the two oxides of phosphorus, P 2 3 and P 2 5 . Arsenic forms with sulphur several compounds, the principal of which are the disulphide, As 2 S 2 , the trisul- phide, As 2 S s , and the pentasulphide, As 2 S 5 . The principal sulphides of antimony are those of the formulas Sb 2 S 3 and Sb 2 S 5 , and of bismuth those of the formulas Bi 2 S 2 and Bi 2 S 3 . Phosphoric Acid, Orthophosphoric Acid, H 3 P0 4 . — The compound to which the name phosphoric acid is generally applied, and from which the best known phosphates are derived, is that which has the formula H 3 P0 4 . To distin- guish it from the other varieties it is called orthophosphoric acid. As has been stated, this is the final product of oxidation of phosphorus in the presence of water. Thus, when phosphorus is boiled with nitric acid it is converted into orthophosphoric acid; and also when phosphorus is burned m the air, and the product dissolved in water, phosphoric acid is formed. In this case the first product of the oxidation is the pentoxide P 2 5 , also known as 324 COLLEGE CHEMISTRY. phosphoric anhydride, and when this is treated with water it is converted into phosphoric acid : P 2 6 + 3H 2 = 2H 3 P0 4 . The occurrence of phosphoric acid in nature has already been referred to in connection with the occurrence of phosphorus, which is found in nature almost exclusively in the form of phosphates, principally as calcium phos- phate, Ca s (POJ 2 , in phosphorite, apatite, and the ashes of tones. In order to prepare the acid two ways suggest them- selves: (1) by oxidizing phosphorus with nitric acid; and (2) by extracting it from one of the natural phosphates, as phosphorite or bone-ash. The first of these methods is better adapted to the preparation of pure phosphoric acid, such as is needed for medicinal purposes ; the latter is used where absolute purity of the product is not required. The preparation of phosphoric acid from a phosphate is not a simple matter. If the acid were volatile or insolu- ble there would be no difficulty in separating it. In the former case it would only be necessary to proceed as in preparing hydrochloric and nitric acids. By adding an acid that is not volatile except at a high temperature, such, for example, as sulphuric acid, and heating, the non-volatile acid drives out the volatile. On the other hand, if phosphoric acid were insoluble in water, it could be separated by adding a soluble acid to one of its soluble salts.. When, for example, nitric acid is added to a solu- tion of potassium tellurite, K 2 TeO s , tellurious acid, being insoluble, is thrown down : K 2 Te0 3 + 2HNO s = 2KNO s + H 2 TeO s . But phosphoric acid is not volatile and is soluble, so that plainly neither of these methods can be used. By treating the calcium salt with sulphuric acid the calcium can be completely separated in the form of calcium sulphate, which is difficultly soluble in water and insoluble in PHOSPHORIC ACID. 3 2 S alcohol. The ideal reaction to be accomplished is that represented in the following equation : £a s (P0 4 ) 2 + 3H 2 S0 4 = 3CaS0 4 + 2H 8 P0 4 . But when sulphuric acid is added to calcium phosphate, only a part of the calcium is thrown down as sulphate, the rest remaining in the form of primary calcium phosphate : Ca s (P0 4 ) 2 + 2H a S0 4 = 2CaS0 4 + CaH 4 (P0 4 ) a . The phosphate thus formed is soluble in water, and the calcium is not easily precipitated from it. By evaporation and addition of sufficient sulphuric acid and alcohol the precipitation can be effected, and a solution of phosphoric acid thus obtained. This acid is not pure, as there are substances in bone-ash which are not removed by the method described. Phosphoric acid for medicinal purposes is often made by dissolving the pentoxide in water. Properties. — When evaporated to the proper consistency the acid forms a thick syrup which slowly solidifies in the form of large crystals. The crystals are deliquescent. When heated to a sufficiently high temperature the acid loses water, as already explained, and yields, first; pyro- phosphonc, and then metaphosphoric acid. It is a tribasic acid, capable of yielding three classes of salts of the (OH (OH (OM general formulas OP \ OH, OP-{ OM, and OP-^ OM, ( OM ( OM ( OM which are known respectively as the primary, secondary, and tertiary phosphates. The primary and secondary phosphates are also known as acid phosphates, and the tertiary salts as neutral or normal phosphates. In these salts it is not necessary that all the hydrogen should be replaced by the same metal. There are salts in which two or three metals take the place of the hydrogen atoms. A phosphate much used in the laboratory, for example, is one in which one hydrogen atom of phosphoric acid is 326 COLLEGE CHEMISTRY. replaced by a sodium atom, and another by the ammonium (OH group, NH,. This salt has the formula OP \ ONa , and ( ONH. is called ammonium sodium phosphate. Another phosphate commonly met with is ammonium magnesium phosphate, I ONH. OP ■< 0. M , which is derived from the acid by replace- ( >JSlg ment of two hydrogen atoms in the molecule by one bivalent magnesium atom, and one by the ammonium group. The changes which these three classes of phosphates undergo when heated are of special interest. The tertiary phosphates are stable. The primary and secondary phos- phates give up all their hydrogen, which passes off in the form of water. Thus, primary sodium phosphate, H 2 NaP0 4 , loses one molecule of water from each molecule of the salt, and is converted into the metaphosphate, NaP0 3 : (OH OP \ OH = O a P(01Sra) + H 2 0. (ONa In general, the primary phosphates are converted into meta- phosphates by heat. "When a secondary phosphate is heated the product is a pyrophosphate, as when secondary sodium phosphate is heated to a sufficiently high temperature it is converted into sodium pyrophosphate : (ONa . ■ gg. OP, roH - t° + H A OP ONa 0POJ» or 2Na 2 HP0 4 = Na 4 P 2 0, + H 2 0, PYROPHOSPHORIC AND METAPHOSPHORIC ACIDS. 327 In general, a secondary phosphate is coverted into a pyrophosphate by heat. Pyrophosphoric Acid, H 4 P a O,. — When phosphoric acid is heated to 200°-300° until a specimen neutralized with ammonia gives a pure white precipitate with silver nitrate, it is completely transformed into pyrophosphoric acid by ' loss of water. The white precipitate referred to is the silver salt of pyrophosphoric acid. The silver salt of orthophosphoric acid is yellow. This difference in color led, many years ago, to a careful investigation of the change in composition which phosphoric acid undergoes when heated, and to the recognition of the existence of pyrophosphoric acid as distinct from orthophosphoric acid; and the study of the relations existing between these acids and metaphosphoric acid has had a strong influence in shaping the views of chemists in regard to the relations between other similar acids. The views at present held in regard to the relations between the common forms of oxygen acids and the so-called normal acids or maximum hydroxides are simply an extension of the ideas first intro- duced into chemistry by Graham in connection with the three varieties of phosphoric acid. The modifications of sulphuric acid seen in the normal acid, S(0H) 6 , the ordi- nary acid, SO,(OH) 2 , and the pyro-acid, H 2 S 2 O t , are examples of the same kind of relations. Pyrophosphates are formed, as we have seen, when the 1 secondary phosphates, like disodium phosphate, HNa 2 P0 4 , are heated. Metaphosphoric Acid, HPO s .— This acid is formed by dissolving phosphorus pentoxide, P 2 5 , in cold water: P,0 6 + H 2 = 2HPO s . It is also formed by heating phosphoric acid to 400° : H a P0 4 = HP0 3 + H 2 0. • Further, the metaphosphates are formed by heating the primary phosphates like primary sodium phosphate, 328 COLLEGE CHEMISTRY. H 2 KaP0 4 . The acid is a vitreous translucent mass, known in the market as glacial phosphoric acid (Acidum phos- phoricum glaciate). It is the more common commercial form of phosphoric acid. It is a monobasic acid, and in composition is analogous to nitric and chloric acids : HPO s Metaphosphoric acid. KN"0, Nitric acid. HC10 3 Chloric acid. When boiled with water in which there is a little nitric acid metaphosphoric acid is readily converted into ortho phosphoric acid: HP0 3 + H 2 = H 3 P0 4 . This transformation is effected also by simply allowing the solution of the meta-acid in water to stand for a time, and by boiling the solution. Phosphorous Acid, H s P0 3 . — This acid is formed when phosphorus trichloride is treated with water. It is also formed together with phosphoric and hypophosphoric acids when phosphorus is allowed to lie in contact with moist air. The acid can be obtained from its solutions by evaporation, when it is deposited in transparent crystals. Phosphorous acid is only dibasic, its salts having the general formula HM 2 P0 3 . This fact has led to the belief that in the acid two of the hydrogen atoms are in com- bination with oxygen in the form of hydroxyl, while the third is in combination with phosphorus as represented in ( H the formula OP \ OH. This conclusion finds further sup- (OH port in the conduct of some derivatives of phosphorous acid. j Hypophosphoric Acid, H 4 P 2 6 , is formed together with 'phosphoric and phosphorous acids when sticks of ordinary phosphorus placed in glass tubes drawn out to a small opening at one end are exposed to the action of moist air. PHOSPHORIC ANHYDRIDE— PHOSPHORUS ANHYDRIDE. 329 Hypophosphorous Acid, H 3 P0 2 , has already been re- ferred to, as its potassium salt is formed in the preparation / / of phosphine by the action of phosphorus upon a solution of potassium hydroxide : -so a = sbao >so » + 2H »°- The product is antimonyl sulphate. The weak basic char- acter of the hydroxides of antimony is shown by the fact that many of its salts are decomposed by water. The salt of antimony which is most commonly met with is the so-called tartar emetic, which appears to be an antimonyl potassium salt of tartaric acid. Tartaric acid is a dibasic acid of the formula C 4 H 4 4 j ^rr, When one of its acid hydrogen atoms is replaced by potassium, and the other / // ANTIMONY COMPOUNDS. 335 by the antimonyl group SbO, the salt thus formed is tartar emetic, C 4 H 4 4 j g| b0 . Antimony tri oxide dissolves in hydrochloric acid, form- ing the trichloride, and this, as has been stated, is decom- posed by water yielding oxychlorides. Antimony Tetroxide, SbjO^ — This compound is most easily obtained by igniting antimonic acid, HjSbO^. Two reactions are of course involved : 2H 8 SbO, = Sb a 5 + 3H a O; Sb 2 6 = Sb,0 4 + 0. It is also formed by igniting the trioxide in the air. At ordinary temperatures the tetroxide is white, but it becomes yellow when heated. Towards strong acids this oxide acts like a weak base. A potassium salt of the formula K 2 Sb 2 6 is known, which is derived from the acid H 2 Sb 2 5 , and this in turn from the simpler acid SbO(OH) 2 by loss of water. The oxide itself is regarded by some as an antimonyl salt of metantimonic acid, Sb0 2 . OH, of the formula Sb0 2 .O.SbO. Antimony Pentoxide, Sb 2 5 . — The tetroxide of antimony does not combine with oxygen to form the pentoxide. The latter can be obtained only by gentle ignition of anti- monic acid, care being taken not to raise the temperature high enough to decompose the pentoxide into the tetroxide and oxygen. The fact that the pentoxide readily yields salts of antimonic acid when treated with basic solutions was mentioned under Antimonic Acid. /Antimony Trisulphide, Sb 2 S 3 . — This compound occurs in nature in considerable quantity and is the chief source of antimony. It is known as stibnite and antimony blende. In some localities, especially in Japan, it occurs in large crystals of great beauty. When heated in the air, or roasted, it is converted into the trioxide, and finally into the tetroxide, while the sulphur escapes as the dioxide. 33 6 COLLEGE CHEMISTRY. Hydrochloric acid, dissolves the trisulphide in the form of the chloride with evolution of hydrogen sulphide : Sh,S 3 + 6H01 = 2SbCl, + 3H„S. Nitric acid converts it into the oxide with separation of sulphur. When a solution of antimony chloride is treated with hydrogen sulphide, the trisulphide is thrown down. This artificially-prepared trisulphide has an orange-red color, while that which occurs in nature is black or gray. The sulphide dissolves in solutions of metallic sulphides, forming salts of sulphantimonious acid, either SbS.SH or Sb(SH) 3 . Antimony Pentasulphide, Sb 2 S 6 , is formed by passing hydrogen sulphide into a solution of antimonic acid or by decomposing a salt of sulphantimonic acid by means of an acid. The action takes place thus : 2H 3 Sb0 4 + 5H a S = Sb 2 S 5 + 8H a O; 2Na 3 SbS 4 + 6H01 = 6Na01 + Sb 2 S 5 + 3H„S. It is, when dry, a golden-yellow powder known as sulphur auratum. It dissolves easily in solutions of metallic sul- phides, forming the sulphantimonates, of which the sodium salt, Na 3 SbS 4 , known as Schlippe's salt, is a good example. The action is represented by this equation : Sb 2 S 5 + 6NaSH = 2Na 3 SbS 4 + 3H 2 S. When heated in the air the pentasulphide gives off enough sulphur to form the trisulphide ; while when the pentoxide is heated it is converted into the tetroxide. The sulphantimonates are decomposed when treated with acids and the pentasulphide is thrown down. ,--' Oxychlorides of Antimony. — Under the head of Anti- mony Trichloride the fact was mentioned that this com- pound is decomposed by cold water as represented in the equation SbCl 3 + H,0 = SbOCl + 2HC1. \ \ OXIDES AND SALTS OF BISMUTH. 337 If, however, hot water is used, the composition of the product approximates to that represented by the formula Sb 4 5 Cl 2 . This complex mixture of oxychlorides is known as the "Powder of Algaroth." It may be regarded as derived from the simple oxychloride by loss of antimony trichloride, thus: 5SbOCl = Sb 4 5 Cl 2 + Sb01 s . Many other oxychlorides besides the two mentioned have been obtained, but they are all more or less closely related to the simple compound SbOCl. Oxides of Bismuth. — The principal compound of bis- muth and oxygen is the trioxide, Bi s O s , which is formed when bismuth is burned in the air. It is a yellow powder. It is basic and forms salts which in composition correspond to the salts of antimony. Like the latter, they are of two classes — the bismuth salts and the bismuthyl salts. The former are derived from the triacid base, Bi(OH) 3 , the latter from the monacid base, BiO(OH). Salts of Bismuth. — The best known salts of bismuth are those which it forms with sulphuric and with nitric acids. There is a sulphate of the formula BiH(S0 4 ) 2 formed by dissolving bismuth oxide in dilute sulphuric acid. The sulphate which is most stable in the presence of water is the bismuthyl salt, (BiO) 2 S0 4 . When bismuth is dissolved in nitric acid and the solution evaporated to dryness the salt Bi(N0 3 ) 3 -j- 10H 2 O is obtained. This salt is decom- posed when heated, and by water, forming basic nitrates of bismuth. The composition of the basic nitrate obtained by decomposing the neutral nitrate with water differs according to the conditions. Hot and cold water produce different results. A solution containing much nitric acid does not give the same result as one which contains little, i etc. As basic bismuth nitrate is used in medicine it is necessary that specific directions should be given for its preparation, in order that a substance of the same com- position should always be obtained. The basic nitrate of * 4, 33 8 COLLEGE CHEMISTRY. bismuth, or the sulnitrate, as it is frequently called in pharmacy, is much used in medicine as a remedy in dysentery and cholera. It is also used as a cosmetic. Bismuth Dioxide, Bi 2 2 , is formed as a brown precipi- tate when potassium hydroxide is added to a solution of bismuth chloride and stannous chloride, SnCl 2 . Bismuth Pentoxide, Bi 2 5 , is formed by oxidizing the trioxide, by means of chlorine, in alkaline solution. Although some experimenters appear to have obtained salts of bismuthic acid, as, for example, KBiO s , others have failed to obtain them. In any case it is evident that the acid properties of the oxide are very weak. Bismuth Trisulphide, Bi 2 S 3 , occurs in nature, and is formed by precipitating bismuth from solutions of its salts with hydrogen sulphide. It dissolves in hot concentrated hydrochloric acid and in nitric acid. It does not dissolve in solutions of the sulphides as the sulphides of arsenic' and antimony do. Bismuth Oxychloride, BiOCl, which in composition is analogous to the simplest form of antimony oxychloride, is thrown down as a white powder when a solution contain- ing bismuth chloride is treated with water : BiCl a + H 2 = BiOCl + 2HC1. $W Family V, Group A. As the members of Group A, Family VII, are related to Group B of the same family; and as the members of Group A, Family VI, are related to the members of Group B of the same family, so the members of Group A, Family V, are related to the members of Group B, which have just been studied. The members of Group A are vanadium, columbium, tantalum, and didymium, all of which are rare. Of these vanadium has been most thoroughly investigated, and columbium next. Vanadium, V (At. Wt. 51.2). — This element occurs in nature in the form of vanadates or salts of vanadic acid, H 3 V0 4 , which is analogous to phosphoric acid. BORON. 339 Vanadic Acid, H s V0 4 , is the most important and best known of the compounds of vanadium. It is the final product of the oxidation of vanadium, and bears to this element the same relation that phosphoric, arsenic, and antimonic acids bear to phosphorus, arsenic, and anti- mony. The vanadates are derived from ortho-, meta-, and pyrovanadic acids, though the most stable ones are the metavanadates, MY0 3 . The free metavanadic acid is known. It is a beautiful golden-yellow compound, which may be used as a substitute for gold bronze. Columbium, Cb (At. Wt. 94). — This element, which is sometimes called niobium, occurs in the mineral colum- bite. Tantalum, Ta (At. Wt. 183). — Tantalum occurs in the minerals columbite and tantalite, accompanied by niobium. Didymium consists of two very similar elements, neody- mium and praseodymium. In some of their compounds they show a resemblance to the members of this group. They form, for example, an oxide of the formula Di 2 5 . On the other hand, they seem to be more closely related to cerium and lanthanum, which are also very rare ele- ments, occurring associated with didymium. These will be further treated of in connection with lanthanum and cerium. Bobon, B (At. Wt. 11). General. — Although the element boron is not a member of the family to which nitrogen and phosphorus belong, it nevertheless resembles the members of this family in some respects. It belongs to the same family as aluminium, and in the composition of its compounds it is undoubtedly similar to aluminium ; but, on the other hand, its oxide is distinctly acidic, while that of aluminium is basic. Occurrence. — Boron occurs in nature chiefly in the form of boric acid, or as salts of this acid, particularly a sodium salt known as borax. Preparation. — From borax and the other borates the acid 340 COLLEGE CHEMISTRY. can easily be obtained. When heated, water is given off, and boron trioxide, B 2 3 , is left : 2B(OH) 3 = B 2 3 + 3H 2 0. By heating the oxide with potassium amorphous boron is obtained. By melting the oxide with aluminium, boron is formed and is dissolved in the molten aluminium, from which, on cooling, it is deposited in crystals. Amorphous boron in almost pure form is obtained by heating borax with magnesium powder. One of the cbief difficulties encountered in preparing boron is to prevent the element from combining with the nitrogen of the air. At the high temperature at which the reduction takes place the two elements combine very readily to form the compound boron nitride, BN . The crystals obtained in the process described are not pure boron, but contain aluminium, or carbon and aluminium, apparently in combination with the boron. The crystals are very hard, and some of them have a high lustre. / Properties. — Amorphous boron is a greenish-brown powder. It burns when heated in the air or in oxygen, | the product being the trioxide B 2 3 . Strong oxidizing j agents, like nitric acid and saltpetre, readily oxidize it, 1 forming boric acid. It combines readily also with many | other elements, as with chlorine, nitrogen, and sulphur. \ When it is brought into the melting hydroxides or car- \bonates of potassium or sodium, it forms borates of the corresponding metals. Boron Trichloride, BC1 3 . — This compound is formed by heating boron in a current of dry chlorine, and by heating a mixture of boron trioxide and charcoal in chlorine: 2B a 3 + 30 + 6C1 2 = 4BC1 3 + 3C0 3 . This reaction is especially interesting on account of its double character. Carbon alone could not reduce the boron trioxide at the temperature employed; nor could the chlorine alone displace the oxygen and form the BORIC ACID. 341 chloride, but when both chlorine and carbon act together these changes take place, one aiding the other. The chloride is a liquid that boils at 17°. Like phos- phorus trichloride, it is easily decomposed by water, forming boric acid, which, as will be seen, is analogous in composition to phosphorous acid and arsenious acid : BC1 S + 3H 2 = B(OH) 3 + 3H01. Boric Acid, B(OH) 3 . — Boric acid occurs free in nature and in the form of salts, of which the principal one is borax. Besides borax, which is a sodium salt derived from tetraboric acid, H 2 B 4 0,, there are other natural borates, as boracite, which is a magnesium salt combined with magnesium chloride; and datholite, which is made up of silicic acid, boric acid, and the element calcium. One of the most interesting natural forms of boric acid is that which is given off from the earth with steam. Such jets of steam are met with in many volcanic regions, and are called fumaroles. In Tuscany many of the fumaroles are charged with small quantities of boric acid, which is somewhat volatile with steam. Those at Monte Cerboli and Monte Kotundo in Tuscany are utilized for the pur- pose of obtaining the boric acid. For this purpose basins are built over the fumaroles and filled with water, so that the steam is condensed and the boric acid dissolved in the water. The solutions formed at the higher levels flow into basins at lower levels, and finally become charged with a considerable quantity of the acid, when it is evaporated to crystallization by the aid of the heat furnished by the f umaroles. The acid obtained in this way is not pure, but it can be purified by recrystallization. Boric acid can also be made from borax by heating the salt in solution with dilute sulphuric acid : Na a B 4 0, + H 2 S0 4 + 5H,0 = Na 2 S0 4 + 4B(OH) 3 . If the solution is sufficiently concentrated the boric acid crystallizes out on cooling. Boric acid is easily soluble in water, and crystallizes 342 COLLEGE CHEMISTRY. from the solution. It is also soluble in alcohol, and this solution burns with a characteristic green flame. The acid is quite volatile with water-vapor. When heated at 100° orthoboric acid loses one molecule of water, and is con- verted into metaboric acid, HB0 2 ; at 160° it yields tetra- boric acid, H 2 B 4 7 ; and at a higher temperature it is converted into boron trioxide or boric anhydride, B 2 3 . These changes are represented in the equations following : H 3 BO s = HB0 2 + H a O; 4H 8 B0 3 = H 2 B 4 7 + 5H a O. The most stable salts are the tetraborates and meta- borates. Borax is the sodium salt of tetraboric acid, Na 2 B 4 0,. The salts of orthoboric acid are unstable. They break down when treated with water, forming. free boric acid and either metaborates or tetraborates. Most of the boric acid obtained from Tuscany is used in the manufacture of borax, a salt which finds extensive application. Salts of Boron. — Although the most characteristic com- pounds of boron are those in which it acts as an acid- forming element, it forms some compounds in which its power as a base-former is shown. Thus, with concen- trated sulphuric acid the trioxide forms a compound which appears to be pyrosulphuric acid, H 2 S a O, , in which the group BO, analogous to antimonyl, SbO, is sub- stituted for one hydrogen. It has the composition (BO)HS 2 0,. Further, when concentrated phosphoric acid acts upon crystallized boric acid, boron phosphate, BP0 4 , is formed. This compound is characterized by great stability. Concentrated acids, for example, do not decom- pose it. It also forms a salt which appears to be analogous to tartar emetic, which, as has been pointed out, is prob- ably antimonyl potassium tartrate, C 4 H 4 6 -j -g . This is the salt represented by the formula 4 H 4 O \ -g- , which may be called boryl potassium tartrate. EXPERIMENTS WITH PHOSPHORIC ACID, ETC. 343 Nitrogen Boride, BN\ — This compound has been referred to in connection with the preparation of boron. It is easily obtained by igniting a mixture of dehydrated borax and ammonium chloride. It forms a white powder, which is insoluble in water, and is characterized by great stability. At red heat it is decomposed by water-vapor into ammonia and boric acid: 2BN + 6H 2 = 2B(OH) s + 2NH 8 . EXPERIMENTS. Phosphoric Acid. Experiment 161. — In a flask connected with an inverted con- denser, as shown in Fig. 63, boil 10 to 15 grams of ordinary phos- Fio. 63. phorus with 250 cc. ordinary commercial nitric acid. If neces- sary, add more acid after a time. Boil gently until the phos- phorus disappears. Evaporate the solution to complete dryness, so as to get rid of all the nitric acid. Experiment 162. — Try the action of a solution of ordinary sodium phosphate on a solution of silver nitrate. Heat a little of the phosphate in a porcelain crucible to redness. After cooling, try the action of the salt left in the crucible on silver nitrate. Arsenic Acid. Experiment 163. — Pass chlorine into water containing arsenic trioxide in suspension, until the oxide is dissolved. Evaporate to 344 COLLEGE CHEMISTRY. crystallization. Into a dilute solution of the product thus ob- tained, to which some hydrochloric acid is added, pass hydrogen sulphide. Explain the changes. E eduction of Arsenic Trioxidb. Experiment 164. — In the bottom of a dry tube of hard glass of the form represented in Fig. 64 put a minute piece of arsenic tri- oxide, and just above it a small bit of charcoal. Heat n gently. Explain the change. Sulphides of Arsenic. Experiment 165. — Pass hydrogen sulphide into a di- lute solution of arsenic trioxide in hydrochloric acid. — Filter off the precipitate, and try the action of ammo- nium sulphide on some of it. Sulphides of Antimony. II Experiment 166. — Pass hydrogen sulphide into a so- il lution of antimonic acid made by treating antimony Fio. 64. with aqua regia and diluting with water. Pass hydro- gen sulphide into a solution of antimony trichloride made by dissolving stibnite or antimony trisulphide in hydrochloric acid. Try the action of ammonium sulphide on the precipitates after filtering. OXYCHLORIDES OE ANTIMONY. ^ Experiment 167. — Treat a solution of antimony trichloride with water. Basic Nitrates of Bismuth. ^ Experiment 168. — Dissolve a little bismuth in nitric acid and evaporate. Add water. Boron. ^Experiment 169. — Make a hot solution of 30 grams crystallized borax in 120 cc. water. Add slowly 10 grams concentrated sul- phuric acid. On cooling, the boric acid will crystallize out. What evidence have you that the substance which crystallizes out of the solution is not borax ? Try the solubility in alcohol of specimens of each. Is there any difference ? Treat a few crys- tals of borax with about 10 cc. alcohol ; pour off the alcohol and set fire to it. Treat a few crystals of the boric acid in the same way. What difference do you observe ? Distil an aqueous solu- tion of boric acid, and determine whether any of the acid passes over with the water-vapor. CHAPTER XIX. CARBON (C, At. Wt. 12) AND ITS SIMPLER COMPOUNDS WITH HYDROGEN AND CHLORINE. Introductory. — Carbon bears to Family IV relations similar to those which nitrogen, oxygen, and fluorine bear to Families V, VI, and VII. Towards hydrogen, as well as towards chlorine and oxygen, carbon is quadrivalent, and towards oxygen it is also bivalent. In this family the maximum oxygen-valence coincides with the hydrogen- valence, while, as has been seen, in Families V, VI, and VII, the oxygen-valence is higher than the hydrogen- valence, the difference becoming greater from Family V to VII. While the higher oxygen compounds of Family IV are acidic, forming acids which are derived from the normal acid, E(OH) 4 , the lower oxides are not generally acid. The hydrogen compounds of the general formula MH 4 , of which there are but two, those of carbon and silicon, have neither acid nor basic properties. Carbon is distinguished by the large number of the compounds into which it enters, all of which are more or less closely related to a comparatively small number of fundamental forms. Silicon also forms a large number of compounds, as we shall see; but these are of a different kind from those obtained from carbon. Occurrence of Carbon. — In general, substances which are obtained from the vegetable or animal kingdom blacken when heated to a sufficiently high temperature, and after- wards, if they are heated in the air, they burn up, as we say. When we consider the great variety of substances found in living things, it appears remarkable that nearly 345 346 COLLEGE CHEMISTRY. all have this property in common. It is due to the fact that nearly all animal and vegetable substances contain the element carbon. When they are heated the other ele- ments present are first driven off in various forms of com- bination, while the carbon is the last to go. Hydrogen and oxygen pass off as water ; hydrogen and nitrogen as ammonia; and much of the carbon also passes off in com- bination with hydrogen, with hydrogen and oxygen, and with nitrogen and hydrogen. If the heating is carried on. in the air, the carbon finally combines with oxygen to form a colorless gas — it burns up. Carbon is the cen- tral element of organic nature. There is not a living thing, from the minutest microscopic animal to the mam- moth, from the moss to the giant tree, which does not contain this element as an essential constituent. The number of the compounds which it forms is almost in- finite, and they present such peculiarities that they are commonly treated of under a separate head, " Organic Chemistry." There is no good reason for this, except the large number of the compounds. For our present purpose, it will suffice to consider the chemistry of the element itself, and of a few of its more important simple com- pounds. From what has already been said, it will be seen that the principal form in which carbon occurs in nature is in combination with other elements. It occurs not only in living things, but in their fossil remains, as in coal. Coal-oil, or petroleum, the formation of which is perhaps due to the action of water on metallic carbides, consists of a large number of compounds which contain only carbon and hydrogen. Most products of plant-life contain the elements carbon, hydrogen, and oxygen. Among the more common of these products may be mentioned sugar, starch, and cellulose. Most products of animal life con- tain carbon, hydrogen, oxygen, and nitrogen. Among them may be mentioned albumen, fibrin, casein, etc. Carbon occurs in the air in the form of carbon dioxide. It also occurs in the form of salts of carbonic acid; the DIAMOND-GRAPHITE. 34? carbonates, which are very widely distributed, forming whole mountain ranges. Limestone, marble, and chalk are varieties of calcium carbonate. Uncombined, the element occurs pure in two very different forms in nature: (1) as diamond; and (2) as graphite, or plumbago. Diamond. — The diamond occurs in few places on the earth, and but little is known as to the conditions which gave rise to its formation. The celebrated diamond beds are in India, Borneo, Brazil, and South Africa. When found, diamonds are generally covered with an opaque layer, which must be removed before its beautiful proper- ties are apparent. The crystals are sometimes regular octahedrons, though usually they are somewhat more complicated, and the faces are frequently curved. It is the hardest substance known. For use as a gem it must be cut and polished. The object in view is to bring out as strikingly as possible its brilliancy by exposing the faces favorably to the action of the light. If heated to a very high temperature without access of air, it swells up and is converted into a black mass resembling coke. The change takes place without loss in weight. Heated to a high temperature in oxygen, it burns up, yielding only carbon dioxide. It is insoluble in all known liquids at ordinary temperatures. It dissolves, however, in molten cast iron and in some other molten metals. Small diamonds have recently (1897) been made by Moissan by dissolving carbon in cast iron with the aid of an electric furnace, and suddenly cooling- the mass. Under these conditions the carbon in the inner parts of the mass is under great pressure while passing into the solid form from solution. Graphite. — Graphite, or plumbago, is found in nature in large quantities. Sometimes it is crystallized, but in forms entirely different from those assumed by the diamond. It can be prepared artificially by dissolving charcoal in molten iron, from which solution, on cooling, graphite is deposited. It can also be prepared by heating 348 COLLEGE CHEMISTRY. charcoal to a high temperature in the electric furnace. This variety differs in some of its properties from that which crystallizes from iron. Both varieties are found in nature. It has a grayish-black color and a metallic lustre. It is quite soft, leaving a leaden-gray mark on paper when drawn across it, and it is hence used in the manufacture of so-called Zea^-pencils. It is sometimes called black-lead. When heated without access of air it remains unchanged. Heated to a very high temperature in the air, or in oxygen, it burns up, forming only carbon dioxide. Like the diamond, it is insoluble in all known liquids at ordinary temperatures. Amorphous Carbon. — All forms of carbon which are not diamond, nor graphite, are included under the name amorphous carbon. The name signifies simply that it is not crystallized. The most common form of amorphous carbon is ordinary charcoal. Charcoal is that form of carbon which is made by the charring process. This consists simply in heating wood without a sufficient supply of air to effect complete com- bustion. The substance almost exclusively used in the manufacture of charcoal is wood. As has already been stated, wood is made up of a large number of substances, nearly all of which, however, consist of the three ele- ments carbon, hydrogen, and oxygen. One of the chief constituents of all kinds of wood is cellulose. Now, when we set fire to a piece of wood, — that is to say, when we heat it up to the temperature at which oxygen begins to act on it, — it burns, if air is present. Under ordinary circumstances the chemical changes that take place are complex; but if care is taken, the combustion can be made complete, when all the carbon is converted into carbon dioxide, and all the hydrogen into water. If, on the other hand, the air is prevented from coming in contact with the wood, as by heating it in a closed vessel, or if it is prevented from coming in contact with it sufficiently to effect complete combustion, the hydrogen is given off partly as water and partly in the form of volatile products AMORPHOUS CARBON. 349 containing carbon and oxygen, as wood spirits or methyl alcohol, pyroligneous or acetic acid, acetone, etc. The carbon, however, is for the most part left behind as char- coal, as there is not enough oxygen to convert it into carbon dioxide. Such a process as that just described, when carried on in closed vessels, is known as destructive distillation or dry distillation. It is also known as the charring process. It is a complex example of a kind of change which we have already had to deal with. When- ever chemical compounds are heated the constituents tend to arrange themselves in forms that are stable at the higher temperature. Sulphites become sulphates; phosphites become phosphates; chlorates become perchlorates; am- monium salts break down into the acids and ammonia; ammonium nitrite is decomposed into nitrogen and water; ammonium nitrate yields nitrous oxide and water; primary phosphates yield metaphosphates; secondary phosphates yield pyrophosphates, etc., etc. Carbon compounds are, in general, more sensitive to the influence of heat than the compounds of other elements, and all are decomposed even at comparatively low temperatures. The above statements will make it possible to under- stand the working of a charcoal-kiln. This consists essentially of a pile of wood arranged to leave spaces between the pieces, and covered with some rough material through which the air will not pass easily, as, for example, a mixture of powdered charcoal, turf, and earth. Small openings are left in this covering, so that after the wood is kindled it will continue to burn slowly. The process is sometimes carried on in structures of brick-work with the necessary number of small openings in the walls. The changes' above mentioned take place, the gases or volatile substances passing out of the top of the kiln, and appear- ing as a dense cloud. In due time the holes through which the air gains access to the wood, thus making the burning possible, are closed, and the burning ceases. Charcoal, which is impure amorphous carbon, is left behind. As wood always contains some incombustible 35° COLLEGE CHEMISTRY. substances in small quantity, these are, of course, found in the charcoal. When the wood or charcoal is burned, these substances remain behind as the ash. Ordinary charcoal is a black, comparatively soft sub- stance. It burns in the air, though not easily, unless the gases that are formed are constantly removed and fresh air is supplied, — conditions which are brought about by a good draught, or by blowing upon the fire with a bellows. It burns readily in oxygen. The product of the combus- tion in the air and in oxygen, when the conditions are favorable, is carbon dioxide, C0 2 . In the air, when the draught is bad, another compound of carbon and oxygen, carbon monoxide, CO, is formed. Heated without access of air, charcoal remains unchanged. Charcoal is insoluble in liquids generally, though it is soluble in molten iron, and it crystallizes from the solution, as we have seen, in the form of graphite, and sometimes as diamond. Coke. — Besides wood charcoal, there are other forms of amorphous carbon, which are manufactured for special purposes, or are formed in processes carried on for the sake of other products. Coke is a form of amorphous carbon which is made by heating ordinary gas-coal without access of air, as is done on the large scale in the manufac- ture of illuminating gas. Coke bears to coal much the same relation that charcoal bears to wood. Lamp-black is a very finely-divided form of charcoal which is deposited on cold objects placed in the flames of burning oils. The oils consist almost exclusively of carbon and hydrogen. When burned in the air they yield carbon dioxide and water. If the flame is cooled down by any means, or if the supply of air is partly cut off, the carbon is not completely burned, the flame " smokes," as we say, and deposits soot. This process is chemically analogous to the deposit of metallic arsenic from a flame of arsine. The soot obtained from the flames of burning oils is made up largely of fine particles of carbon, though some of the unchanged oils are contained in it. It is used in the manufacture of printing-ink. As carbon is acted upon AMORPHOUS CARBON. 351 directly by Tery few substances, and is not soluble, it is almost impossible to destroy the color of printing-ink without destroying the material upon which it is im- pressed. Bone-Mack, or Animal Charcoal, is a form of amorphous carbon that is made by charring bones. Bones consist of about one-third organic matter and two-thirds incom- bustible matter, mostly calcium phosphate. When charred, the organic matter undergoes the changes briefly described under the head of Charcoal, while the incombustible con- stituents remain unchanged. As the organic matter is distributed through the substance of the bones the char- coal obtained in this way is in a very fine state of division, but it is mixed with several times its own weight of min- eral matter. In order to remove the latter the bone-black must be treated with an acid, as hydrochloric acid, and afterwards thoroughly washed with water. .An efficient variety of animal charcoal is made, further, by mixing blood with sodium carbonate, charring, and afterwards dissolving out the sodium carbonate with water. Bone-black and wood-charcoal are very porous, and have the power to absorb gases. When placed in air containing bad-smelling gases these are absorbed, and the air is thus to some extent purified. When water contain- ing disagreeable substances is treated with charcoal, these are wholly or partly absorbed, and the water is improved. Charcoal-filters are therefore extensively used. A char- coal-filter to be efficient should be of good size, and from time to time the charcoal should be taken out and renewed. The small filters which are screwed into faucets are of little value, as the charcoal soon becomes charged with the objectionable material which is present in the water, and is then a source of contamination rather than a means of purification. The power of charcoal to absorb gases depends upon its porosity. That from some varieties of wood is more porous than that from other varieties. Box- wood charcoal has been shown to absorb 90 times its own volume of ammonia gas, 35 times its volume of carbon 35 2 COLLEGE CHEMISTRY. dioxide, and nearly twice its volume of hydrogen. Char- coal from cocoa-nut wood absorbs 172 times its volume of ammonia, and 68 times its volume of carbon dioxide. Some coloring matters can be removed from liquids by passing the - liquids through bone-black filters. On the large scale, this fact is taken advantage of in the refining of sugar. The solution of sugar first obtained from the cane or beet is highly colored; and, if it were evaporated, the sugar deposited from it would be dark-colored. If, however, the solution is first passed through bone-black niters, the color is removed, and now, on evaporating, white sugar is deposited. In the laboratory constant use is made of this method of decolorizing liquids. The action can easily be shown by adding a little bone-black to a solution containing some litmus or indigo. If the solution is digested for a short time with the bone-black, and then passed through a filter, it will be found that the coloring matter is removed. Charcoal does not undergo decay in the air or under water nearly as readily as wood. That is another way of stating the chemical fact that the substances of which wood is made up are more susceptible to the action of other chemical substances than charcoal is. "We have one good illustration of this, indeed, in the relative ease with which charcoal and Wood burn in the air. Piles that are driven below the surface of water are sometimes charred to protect them from the action of those substances which cause decay. Coal. — Under this head are included a great many kinds of impure amorphous carbon which occur in nature. Although we might distinguish between an almost infinite number of kinds of coal, for ordinary purposes they are divided into hard and soft coals, or anthracite and bitumin- ous coals. Then there are substances more nearly allied to wood called lignite, and those which represent a very early stage in the process of coal-formatjon, viz., peat. A close examination of all these varieties has shown that they have been formed by the gradual decomposition of DIAMOND, GRAPHITE, AND CHARCOAL. 353 vegetable matter in an insufficient supply of air. The process has been going on for ages. Sometimes the sub- stances have, at the same time, been subjected to great pressure, as can be seen from the position in which they occur in the earth. The products in the earlier stages of the coal-forming process are more closely allied to wood than those in the later stages. All forms of coal contain other substances in addition to the carbon. The soft coals are particularly rich in other substances. When heated they give off a mixture of gases and the vapors of volatile liquids. The gases are, for the most part, useful for illuminating purposes. The liquids form a black, tarry mass known as coal tar, from which many valuable compounds of carbon are obtained. The gases are passed through water for the purpose of removing certain impuri- ties. This water absorbs ammonia, and forms the am- moniacal liquor of the gas-works, which, as has been stated, is the principal source of ammonia. Diamond, Graphite, and Charcoal are Different Forms of the Element Carbon "We have seen that oxygen presents itself in two forms — ordinary oxygen and ozone. Ozone is made from oxygen, and oxygen from ozone, with- out any increase or decrease in weight; and in general the compounds obtained by the combination of other elements with oxygen are identical with those obtained by the com- bination of the same elements with ozone. So, also, there are several varieties of sulphur, two of which crystallize in different forms. There are, further, three or four differ- ent modifications of the element phosphorus, and these differ from one another in a very marked way. The ex- planation of the difference between oxygen and ozone is that the molecule of the former is made up of two atoms, while that of ozone is made up of three, which are in a state of unstable equilibrium. This explanation is reached through a study of the specific gravity of the two gases. At present no satisfactory explanation can be given of the difference between the varieties of phosphorus and between the varieties of sulphur. It will probably be shown to be 354 COLLEGE CHEMISTRY. due to the way in which the atoms are grouped together in the molecules, and also to the way in which the mole- cules are grouped together to form the masses. Carbon, as we have seen, occurs in three distinct forms. It is difficult to conceive that the black, porous charcoal, and the dull, gray, soft graphite are chemically identical with the hard, transparent, brilliant diamond. Yet this is undoubtedly the case, as can be shown by a very simple experiment. Each of the substances when burned in oxygen yields carbon dioxide. Now, the composition of carbon dioxide is known, so that, if the weight of the carbon dioxide formed in a given experiment is known, the weight of the carbon in it is also known. When a gram of pure charcoal is burned it yields 3| grams carbon dioxide, and in this quantity of carbon dioxide there is contained exactly one gram of carbon. Further, when a gram of graphite is burned the same weight (3| grams) of carbon dioxide is obtained as in the case of charcoal; and the same thing is true of diamond. It follows from these facts that the three forms of matter known as charcoal, graphite, and diamond consist only of the element carbon. The explanation of the difference is not known, but, as in the cases of phosphorus and sulphur, it will probably be found to be in the different ways in which the atoms are arranged in the molecule, and the molecules in the masses. Notwithstanding the marked differences in their appear- ance and in many of their physical properties, the three forms of carbon have, as we have seen, some properties in common. They are insoluble in all known liquids at ordinary temperatures. They are tasteless, inodorous, and infusible. When heated without access of air they remain unchanged, unless the temperature is very high, when the diamond swells up and is converted into a mass resembling coke — a change which is connected with a rearrangement of the particles in an irregular way, so that the substances cease to be crystalline, or become amorphous. „ CHEMICAL CONDUCT OF CARBON. 355 Chemical Conduct of Carbon At ordinary temperatures carbon is an inactive element. If it is left in contact with any one of the elements, no chemical change takes place. It will not combine with any of them unless the tempera- ture is raised. At higher temperatures, however, it com- bines with several of them with great ease, especially with oxygen. Under proper conditions it combines also with nitrogen, with hydrogen, with sulphur, and with many other elements. It combines with oxygen either directly, as when it burns in the air or in oxygen; or it abstracts oxygen from some of the oxides. The direct combination of oxygen and carbon has already been seen in the burn- ing of charcoal in oxygen, and is familiar to every one in the fire in a charcoal furnace. That carbon dioxide is the product formed can be shown by passing the gas through lime-water or baryta-water, when insoluble calcium or barium- carbonate will be thrown down. The reason why lime-water or baryta- water is used is simply that an insolu- ble compound is formed, and this can be seen, and it can be separated from the liquid and examined. The reaction which takes place is represented thus : Ca(OH) a + C0 2 = CaCO, + H,0; Calcium Carbon Calcium hydroxide dioxide carbonate Ba(OH) J + C0 2 = BaCO s + H 2 0. Barium Carbon Barium hydroxide dioxide carbonate No other common gas acts in this way on these solutions. Hence, when, under ordinary circumstances, a gas is passed into lime-water and an insoluble compound is formed, we may conclude that the gas is carbon dioxide, though this conclusion may require further proof. The abstraction of oxygen from compounds by means of carbon may be illustrated in a number of ways. Thus, when powdered copper oxide, CuO, is mixed with powdered charcoal, and the mixture heated in a tube, carbon dioxide is given off, and can be detected as in the last experiment mentioned. Copper is left behind, and, if the proportions are properly selected, all the carbon will 35 6 COLLEGE CHEMISTRY. pass off as carbon dioxide, and only the copper be left behind : 2CuO + = 2Cu + C0 2 . In a similar way, arsenious oxide, As 2 3 , gives up its oxygen to carbon. This fact furnishes indeed a delicate method for the detection of the substance. If a little is placed in the bottom of a small tube, and above it a small- piece of charcoal, then when heat is applied the arsenious oxide sublimes, and as its vapor passes the heated charcoal the oxygen is abstracted, and the element arsenic, being also somewhat volatile, is deposited just above the charcoal in the form of a lustrous mirror on the walls of the tube. The reaction is 2As 2 3 + 30 = 4As + 3C0 2 . As has already been explained, the abstraction of oxygen from a compound is known as reduction. Hence, carbon is called a reducing agent. It is indeed the reducing agent which is most extensively used in the arts. Its chief use is in extracting metals from their ores, which are the forms in which they occur in nature. Thus, iron does not occur in nature as iron, but in combination with other elements, especially with oxygen. In order to get the metal the ore must be reduced, or, in other words, the oxygen must be extracted. This is invariably accomplished by heating it with some form of carbon, either coke or charcoal. The elements chlorine, oxygen, nitrogen, and hydrogen being gases, and the products formed when the first three combine with hydrogen being also gaseous or convertible into vapor, it is a comparatively easy matter to study the relations between the volumes of the combining gases and the volumes of the products formed. It is, however, im- possible to determine the ratio between the volume of carbon gas and that of other gases with which it combines. Compounds of Carbon with Hydrogen, or Hydrocarbons —Carbon forms a large number of compounds with hydrogen. These so-called hydrocarbons, among which EXPERIMENTS WITH CARBON. 357 may be mentioned marsh-gas or methane, CH 4 , olefiant gas or ethylene, C 2 H 4 , acetylene, 2 H a , and benzene, C 6 H 6 , will be treated of in a later chapter together with some other typical compounds of carbon. EXPERIMENTS. v Carbon. — Bone-black Filters. Experiment 170. — Make a filter of bone-black by fitting a paper filter into a funnel 12 to 15 mm. (5 to 6 inches) in diameter at its mouth. Half fill this with bone-black. Pour a dilute solution of indigo through the filter. If the conditions are right the solu- tion will pass through colorless. Do the same thing with a dilute solution of litmus. If the color is not completely removed by one filtration, heat and filter again. The color can also be removed from solutions by putting some bone-black into them and boiling for a time. Try this with half a litre each of the litmus and indigo solutions used in the first part of the experiment. Use about 4 i o 5 grams bone-black in each case. Shake the solution frequently while heating. Charcoal absorbs Gases. Experiment 171. — Collect over mercury in glass tubes some ammonia gas, and some carbon dioxide. Introduce into each Fig. G5. a piece of charcoal, which has been heated in a Bunsen-burnet flame in order to drive out gases which may be contained in the pores. 35 8 college chemistry. Caebon combines with Oxygen to fobm Caebon Dioxide. Experiment 172. — Put a small piece of charcoal in a piece of hard-glass tube. Heat the tube, and pass oxygen through it. Pass the gases into clear lime-water. Arrange the apparatus as shown in Fig. 65. A is a large bottle containing oxygen ; B is a cylinder contain- ing sulphuric acid ; C is a U-tube containing calcium chloride ; D is the hard-glass tube containing the charcoal ; E is the cylin- der with clear lime-water. Explain all that takes place. Caebon seduces some Oxides when heated with them. Experiment 173. — In a small hard-glass tube closed at one end put a little of a mixture of about equal parts of powdered copper oxide and powdered charcoal. Attach a short piece of rubber tubing to the open end of the glass tube, as shown in Fig. 66. Heat the tube and pass the gas which is given off into lime-water contained in a small test-tube. Is it carbon dioxide ? What evidence have you that oxygen has been extracted from the copper oxide? Compare the substance left in the tube with metallic copper. Treat both with nitric acid, with sulphuric acid. ^ Experiment 174. — Repeat Experiment 165 with somewhat larger quantities of the substances, and examine the gas given off. Hydbocaebons. "^ Experiment 175. — Make marsh-gas by heating in a retort a mixture of 20 grams sodium acetate, 20 grams potassium hydrox- ide, and 30 grams slaked lime. Collect some of the gas over water. Is it a combustible gas ? Experiment 176. — Make ethylene as follows : In a flask of 2 to 3 litres capacity put a mixture of 25 grams alcohol and 150 grams ordinary concentrated sulphuric acid. Heat to 160° to 170°, and add gradually through a funnel-tube about 500 cc. of a mixture of 1 part of alcohol and 2 parts of concentrated sulphuric acid. Pass the gas through three wash bottles containing, in order, con- centrated sulphuric acid, caustic soda, and concentrated sul- phuric acid. Collect some of the gas over water. la it com- bustible ? OHAPTEE XX. SIMPLER COMPOUNDS OF CARBON WITH OXYGEN, AND WITH OXYGEN AND HYDROGEN. General. — The final product of oxidation of carbon is carbon dioxide, and the final product of reduction is marsh-gas, but between these two limits there are a num- ber of interesting derivatives, just as there are a number of compounds of sulphur between hydrogen sulphide and sulphuric acid; a number of compounds of nitrogen between ammonia and nitric acid; and a number of com- pounds of phosphorus between phosphine and phosphoric acid. Some of these will be treated of in a later chapter. Carbon Dioxide, C0 2 . — The principal compound of carbon and oxygen is carbon dioxide, C0 2 , commonly known as carbonic acid gas. Under the head of The Air attention was called to the fact that this gas is a constant constituent of the air, though its relative quantity is small — about 3 parts in 10,000. It issues from the earth in many places, particularly in the neighborhood of volcanoes. Many mineral waters contain it in large quantity, promi- nent among which are the waters of Pyrmont, Selters, and the Geyser Spring of Saratoga. In small quantity it • is present in all natural waters. In combination with bases it occurs in enormous quantities, particularly in the form of calcium carbonate, CaC0 3 , varieties of which are ordinary limestone, chalk, marble, and calc-spar. Dolo- mite, which forms mountain-ranges, being particularly 35q 3 6 ° COLLEGE CHEMISTRY. abundant in the Swiss Alps, is a compound containing calcium carbonate and magnesium carbonate, MgC0 3 . Carbon dioxide is constantly formed in many natural processes. Thus, all animals that breathe in the air give off carbon dioxide from the lungs. That the gases from the lungs contain carbon dioxide can easily be shown by passing them through lime-water, when a precipitate of calcium carbonate is formed. That carbon dioxide is formed in the combustion of charcoal and wood has already been shown. In a similar way it can be shown that the gas is formed whenever any of our ordinary combustible substances are burned. From our fires, as from our lungs, and from the lungs of all animals, then, carbon dioxide is constantly given off. Further, the natural processes of decay of both vegetable and animal matter tend to convert the carbon of this matter into carbon dioxide, which then finds its way prin- cipally into the air. The process of alcoholic fermenta- tion, and some other similar processes, also give rise to the formation of carbon dioxide. In all fruit-juices there is contained sugar. When the fruits ripen, fall to the earth, and undergo spontaneous change, the sugar is converted into alcohol and carbon dioxide. We see, thus, that there are many important sources of supply of carbon dioxide, and it will be readily understood why the gas should be found everywhere in the air. Preparation. — The easiest way to get carbon dioxide not mixed with other substances is by adding an acid to a salt of carbonic acid or a carbonate. In the decomposi- tion of the carbonates by other acids we see exemplified the same principle as that which is involved in setting nitric acid free from a nitrate, or hydrochloric acid from sodium chloride by sulphuric acid, and more particularly in the liberation of sulphur dioxide from a sulphite. In all these cases the products are volatile, and therefore, when a fion-volatile acid is added to the salts, decomposi- tion takes place. Sulphites do not yield the corresponding acid, but this breaks down into water and the anhydride: CARBON DIOXIDE. 361 ( 2NaNO s + H 2 SO. = Na.SO. + 2rTN0 3 ; ( 2NaCl + H 2 S0 4 = Na 2 S0 4 + 2HC1. ( Na 8 SO s + H 2 S0 4 = Na 2 S0 4 + H 2 SO s ; \ H 2 S0 3 = S0 2 + H 2 0. ( Na 2 C0 3 + H 2 S0 4 = Na 2 S0 4 + H 2 C0 3 ; |lI 2 C0 3 =C0 2 + H.0. Any acid that is not volatile at the ordinary temperature will decompose a carbonate and cause an evolution of carbon dioxide. The action between sodium carbonate and hydrochloric acid is represented in this way: Na 2 C0 3 + 2HC1 = 2NaCl + C0 2 + H 2 0; that between nitric acid and sodium carbonate in this way: Na 3 C0 5 + 2HNO, = 2NaNO s + C0 2 + H 2 0. For the purpose of preparing carbon dioxide in the laboratory, calcium carbonate, in the form of marble or limestone, and hydrochloric acid are commonly used. The reaction involved is represented thus : CaOO, + 2HC1 = CaCl a + C0 2 +H 2 0. The apparatus used is the same as that used in making hydrogen from zinc and sulphuric acid. As the gas is somewhat soluble in water, it is best for ordinary purposes to collect it by displacement of air, the vessel being placed with the mouth upward, as the gas is considerably heavier than air. Properties. — Carbon dioxide is a colorless gas at ordi- nary temperatures. When subjected to a low temperature and high pressure it is converted into a liquid. Liquid carbon dioxide is now manufactured on the large scale for use as a fire-extinguisher, and for the purpose of charging liquids with the gas. When some of the liquid is exposed to the air evaporation takes place so rapidly that a great deal of heat is absorbed, and some of the liquid becomes solid. The gas has a slightly acid taste and smell. It is not combustible, nor does it support combustion. It is 362 COLLEGE CHEMISTRY not combustible for the same reason that water is not: because it already holds in combination all the oxygen it has the power to combine with. Before it can burn again it must first be decomposed. As regards the statement that it does not support combustion, it should be remarked that this is only relatively true. The compound does not easily give up oxygen, but to some substances it does give it up, and some such substances burn in it. For example, the element potassium, which, as we have seen, has the power to decompose water, has also the power to decom- pose carbon dioxide if heated in it to a sufficiently high temperature, and when the decomposition once begins, it proceeds with brilliancy, the act being accompanied by a marked evolution of heat and light. Carbon dioxide is much heavier than air, its specific gravity being 1.529. A litre of the gas under standard conditions of temperature and pressure weighs 1.977 grams. It dissolves in water, one volume of water dissolving about one volume of the gas at the ordinary temperature. As is the case with all gases, when the pressure is increased the water dissolves more gas, and when the pressure is removed the gas again escapes. The so-called "soda-water" is simply water charged with carbon dioxide under pressure. The escape of the gas when the water is drawn is familiar to every one. The name soda-water has its origin in the fact that the carbon dioxide used in charging the water is frequently made from primary or acid sodium carbonate, NaHC0 3 , which is also called soda or bicarbonate of soda. Relations of Carbon Dioxide to Chemical Energy. — Carbon has the power to combine with oxygen, and in so doing a definite quantity of heat is evolved. A kilogram of carbon represents a certain quantity of chemical energy, which we can get from it first in the form of heat, and by transformation, in other forms of energy, as motion, elec- trical energy, etc. After the kilogram of carbon has been burned it no longer represents the energy it did in the form of carbon. A body of water elevated ten or fifteen feet represents a certain quantity of energy which can be RESPIRATION. 363 obtained by allowing the water to fall upon the paddles of a water-wheel connected with the machinery of a mill. After the water has fallen, however, it no longer has the power to do work, or it has none of the energy which it possessed by virtue of its position. In order that it may again do work it must again be lifted. So, too, in order that the carbon in carbon dioxide may again do work the compound must be decomposed. Respiration. — It was stated above that carbon dioxide is given off from the lungs just as it is from a fire. It is a waste-product of the processes taking place in the animal body. Just as it cannot support combustion, so also it cannot support respiration. It is not poisonous any more than water is; but it cannot supply the oxygen which is needed for breathing purposes, and hence animals die when placed in it. They die by suffocation, very much as they do in drowning. Any considerable increase in the' quantity of carbon dioxide in the air above that which is normally present is objectionable, for the reason that it decreases the proportion of oxygen in the air which is breathed. If, however, pure carbon dioxide is introduced into the air, it has been found that as much as 5 per cent may be present without serious results to those who breathe it. In a badly-ventilated room in which a number of people are collected, and lights are burning, it is well known that in a short time the air becomes foul, and bad effects, such as headache, drowsiness, etc., are produced on the occupants of the room. These effects appear to be due, not to the carbon dioxide, but largely to other waste- products which are given off from the lungs in the process of breathing. The gases given off from the lungs consist of nitrogen, oxygen, carbon dioxide, and water-vapor. Besides these, however, there are many substances in small quantity, in a finely-divided condition, which contain carbon, and are in a state of decomposition. These act as poisons, and they are the chief cause of the bad effects experienced in breathing air which is contaminated by the exhalations from the lungs. As carbon dioxide is given 3^4 COLLEGE CHEMISTRY. off from the lungs at the same time, the quantity of this gas present is^ roughly proportional to the quantity of the organic impurities. Hence, by determining the quantity of carbon dioxide it is possible to form an opinion as to whether the air of a room occupied by human beings is fit for use or not. As carbon dioxide is formed in the earth wherever an acid solution comes in contact with a carbonate, the gas is frequently given off from fissures in the earth. It is hence not infrequently found in old wells which have not been in use for some time, and deaths have been caused by descending into these wells for the purpose of repairing them. The gas is also frequently met with in mines, and is called choke-damp by the miners. The miners are aware that after an explosion caused by fire-damp there is danger of death from choke-damp. The reason of the presence of this gas after an explosion is clear. When fire-damp, or marsh-gas, explodes with air the carbon is oxidized to choke-damp, or carbon dioxide, and the hydrogen to water. Air in which a candle will not burn is not fit for breathing purposes. Carbon Dioxide and Life. — The rSle played by carbon dioxide in nature is extremely important and interesting. The carbon contained in living things is obtained from carbon dioxide, and generally returns to this form when life ceases. We have seen that all living things contain carbon as an essential constituent. Whence comes this carbon ? Animals eat either the products of plant-life or other animals which derive their sustenance from the vegetable kingdom. The food of animals comes, then, either directly or indirectly from plants. But plants derive their sustenance largely from the carbon dioxide of the air. The plants have the power to decompose the gas with the aid of the direct light of the sun, and they then build up the complex compounds of carbon which form their tissues, using for this purpose the carbon of the earbon dioxide which they decompose. Many of these compounds are fit for food for animals; that is to say, they ENERGY STORED UP IN PLANTS. 365 are of such composition that the forces at work in the animal body are capable of transforming them into animal tissues, or of oxidizing them, and thus keeping the tem- perature of the body up to the necessary point. That part of the food which undergoes oxidation in the body plays the same part as fuel in a stove. It is burned up with an evolution of heat, the carbon being converted into carbon dioxide, which is given off from the lungs. From fires and from living animals carbon dioxide is returned to the air, where it again serves as food for plants. When the life-process stops in the animal or the plant, decomposition begins; and the final result of this, under ordinary circum- stances, is the conversion of the carbon into the dioxide. Energy Stored up in Plants. — It will thus be seen that under the influence of life and sunlight carbon dioxide is constantly being converted into compounds containing carbon which are stored up in the plants. These com- pounds are capable of burning, and thus giving heat; or some of them may be used as food by animals, when they assume other forms under the influence of the life-process of the animals. As long as life continues, plants and animals are storehouses of energy. When death occurs, the carbon compounds begin to pass back to the form of carbon dioxide, and the chemical energy is transformed partly into heat, and is thus, as we say, dissipated. The power to do work, which the carbon compounds of plants and animals possess, comes from the heat of the sun. It takes a certain quantity of this heat, operating under proper conditions, to decompose a certain quantity of carbon dioxide, and elaborate the compounds contained in the plants. When these compounds are burned they give out the heat which was absorbed in their formation during the growth of the plants. These compounds are said to possess chemical energy. This has its origin in heat, and is capable of reconversion into heat. The transformation of the energy of the sun's heat into chemical energy lies at the foundation of all life. As the heat of the sun acting upon the great bodies of water and on the air gives rise to 366 COLLEGE CHEMISTRY. the movements of water which are so essential to the exist- ence of the world as it is, so the action of the- sun's rays on carbon dioxide, under the influence of the delicate and inexplicable mechanism oi the leaf of the plant, gives rise to those changes in the forms of combination of the ele- ment carbon which accompany and are fundamental to the wonderful process of life. Carbonic Acid and Carbonates. — When carbon dioxide is passed into water the solution has a slightly acid reac- tion. The solution will act upon bases and form salts. The formula of the sodium salt formed in this way has been shown to be Na 2 C0 3 ; that of the potassium salt, K 2 C0 3 ; etc. These salts are plainly derived from an acid, H 2 00, , which is called carbonic acid. It is probable that this acid is contained in the solution of carbon dioxide in water. It is, however, so unstable that it breaks down into carbon dioxide and water: H 2 C0 3 = C0 2 + H a O. The formation of a salt by the action of carbon dioxide on a base takes place as shown in the following equations : 2KOH +C0 2 = K 2 CO s + H 2 0; Ca(OH) 2 + 00 2 = CaCO s + H 2 0. With the acid the action would take place as represented thus: 2KOH + H 2 C0 3 = K 2 CO a + 2H 2 0; Oa(OH) 2 + H 2 00 s = CaC0 3 + 2H 2 0. There is perfect analogy between the action of carbon dioxide and that of sulphur dioxide on basic solutions. With potassium hydroxide and calcium hydroxide, sulphur dioxide acts as represented in the following equations: 2KOH + S0 2 = K 2 S0 3 + H 2 0; Ca(OH) 2 + S0 2 = OaSO s + H a O. CARBON MONOXIDE. 367 The products formed are sulphites or salts of sulphurous acid. Like sulphurous acid, carbonic acid is dibasic, and forms two series of salts, the primary and secondary, or the acid and normal salts. The primary or acid salts have the general formula HMC0 3 , and the secondary or normal salts have the general formula M 2 C0 3 . Examples of the former are HKC0 3 , HNaC0 3 , CaH 2 (C0 3 ) 2 , etc.; and of the latter K 2 C0 3 , Na 2 C0 3 , CaC0 3 , BaC0 3 , etc. The secondary or normal salts which carbonic acid forms with the most strongly marked metallic elements, viz., potassium and sodium, are not decompose'd by heat, but all other carbonates are decomposed by heat more or less easily, according to the strength of the base. Calcium carbonate when ignited loses carbon dioxide, and lime, or calcium oxide, remains behind : CaCO s = CaO + C0 2 . Carbon Monoxide, CO. — When a substance containing carbon burns in an insufficient supply of air, — as, for example, when the draught in a furnace is not strong enough to remove the products of combustion and supply fresh air, — the oxidation of the carbon is not complete, and the product, instead of being carbon dioxide, is carbon monoxide, CO. This compound can also be made by extracting oxygen from carbon dioxide. It is only neces- sary to pass the dioxide over heated carbon, when reaction takes place as represented thus : CO a + C = 2CO. This method of formation is illustrated in coal-fires, and can be well observed in an open grate. The air has free access to the coal, and at the surface complete oxidation takes place. But that part of the carbon dioxide which is formed at the lower part of the grate is drawn up through the heated coal, and is partly reduced to carbon monoxide. When the monoxide escapes from the upper part of the grate it again combines with oxygen, or burns, 368 COLLEGE CHEMISTRY. giving rise to the characteristic blue flame always noticed above a mass of burning anthracite coal. The monoxide is also formed by passing steam over highly-heated carbon, when this reaction takes place : C -)- H 2 = CO + H 2 . This is the reaction made use of in the manufacture of " water-gas." The gas thus obtained is largely a mixture of hydrogen and carbon monoxide. The gas is enriched by passing it through a furnace in which it is mixed with highly-heated . vapors of hydrocarbons from petroleum. The main reaction, the decomposition of water by heated carbon, is effected in large furnaces filled with anthracite coal. The coal is first heated to a high temperature by setting fire to it, the products of combustion being allowed to escape. When it is hot enough, the air is shut off and steam passed rapidly in, when the decomposi- tion of the water by the carbon takes place. Soon the mass becomes so much cooled that the reaction stops. The steam is then cut off and air turned on again, and so on. The easiest way to make carbon monoxide is by heating oxalic acid, which is a compound of carbon, hydrogen, and oxygen, of the formula C 2 H 2 4 , with five to six times its weight of concentrated sulphuric acid. The change which takes place is represented thus : C 2 H 2 4 = C0 2 + CO + H 2 0. Both the dioxide and monoxide of carbon are formed. Both are gases. In order to separate them the mixture is passed through a solution of sodium hydroxide, which takes up the carbon dioxide, forming sodium carbonate, and allows the monoxide to pass. Carbon monoxide is a colorless, tasteless, inodorous gas, insoluble in water. It burns with a pale-blue flame, forming carbon dioxide. It is exceedingly poisonous when inhaled. Hence it is very important that it should not be EXPERIMENTS WITH CARBON DIOXIDE. 369 allowed to escape into rooms occupied by human beings. We not infrequently hear of deaths caused by the gases from coal-stoves. The most dangerous of the gases given off from these stoves is probably carbon monoxide. A pan of smouldering charcoal gives off this gas, and the fact that it is poisonous is well known. It has been used to a considerable extent for the purpose of suicide. The poisonous character of carbon monoxide has led to a great deal of discussion and to some legislation on the subject of "water-gas." The question has been repeatedly raised whether the government should allow the manufacture of the gas. There is no doubt of the fact that it is a danger- ous substance, and that it should not be allowed to escape into the air is obvious. Wherever it is used special precau- tions should be taken to guard against leaking. There is no doubt that it is more poisonous than coal-gas. At high temperatures carbon monoxide has a very strong tendency to combine with oxygen, and is hence a good reducing agent. In the reduction of iron from its ores, the carbon monoxide formed in the blast-furnace plays an important part in the reducing process. At ordinary temperatures the gas does not combine readily with oxygen. When passed over some substances which are rich in oxygen, as, for example, chromic anhydride, 0r0 3 , and potassium permanganate, KMn0 4 , in acid solution, it takes up oxygen even at the ordinary temperature. It unites with chlorine in the direct sunlight, and forms the compound known as carbonyl chloride or phosgene, C001 a . EXPEKIMENTS. Carbon Dioxide is formed when a Carbonate is treated with an acid. Experiment 177.— In test-tubes add successively dilute hy- drochloric, sulphuric, nitric, and acetic acids to a little sodium carbonate. In each case pass the gas given off through lime- water, and insert a burning stick in the upper part of each tube. Perform the same experiments with small pieces of marble. 37° COLLEGE CHEMISTRY. Preparation and Properties of Carbon Dioxide. Experiment 178. — Arrange an apparatus as shown in Fig. 67. In the flask put some pieces of marble or lime- stone, and pour ordinary hydrochloric acid on it. The gas should be collected by dis- placement of air, the vessel being placed with the mouth upward. Collect several cylinders or bottles full of the gas. Into one introduce successively a lighted candle, a burning stick, a bit of burning phosphorus. With another proceed as if pouring water from it. Pour the invisible gas upon the flame of a burning candle. Pour some of the gas from one vessel to another, and show that it has been transferred. Balance a beaker on a good-sized pair of scales, and pour carbon dioxide into it. If the balance is at all sensitive, the pan on which the Fig. 67. beaker is placed will sink. Carbon Dioxide is given ope from the Lungs. ^ Experiment 179. — Force the gases from the lungs through some lime-water by means of an apparatus arranged as shown in Fig. 68. Formation of Carbonates. S*. Experiment 180. — Pass carbon di- oxide into a solution of potassium hydroxide to saturation. Deter- mine whether a carbonate is in solu- . tion or not. X Experiment 181. — Pass carbon di- oxide into 50 to 100 cc. clear lime- water. Filter off the white insoluble substance. Try the action of a little acid on it. What evidence have you that it is a carbonate ? \_ Experiment 182. — Pass carbon dioxide first through a little water to wash it, and then into 50 to 100 cc. clear dilute lime- water. Continue to pass the gas for some time after the pre- cipitate is formed. The precipitate dissolves. Heat the solution. What happens ? Explain these reactions. Fiq. 68. EXPERIMENTS WITH CARBON MONOXIDE. 371 I Preparation and Properties of Carbon Monoxide. ■^ Experiment 183. — Put 10 grams crystallized oxalic acid and 50 to 60 grams concentrated sulphuric acid in an appropriate flask. Connect with two WoulfE's flasks containing a solution of caustic soda. Heat the contents of the flask gently. Collect some of the gas over water. Set fire to some, and notice the characteristic blue flame. Carbon Monoxide is a good Reducing Agent. Experiment 184. — Pass carbon monoxide over some heated copper oxide contained in a hard-glass tube. Is the oxide re- duced ? How do you know ? Is carbon dioxide formed ? What evidence have you ? "Was the carbon monoxide used free of car- carbon dioxide ? If not, what evidence have you that carbon dioxide is formed in this experiment ? Experiment 185.— Pass carbon dioxide over heated charcoal in a hard-glass tube. What is formed ? CHAPTER XXI. ILLUMINATION.— FLAME.— BLOWPIPE. COMPOUNDS OF CARBON WITH NITROGEN AND SULPHUR. Introduction. — As the substances used for illumination contain carbon, and the chemical processes involved con- sist largely in the oxidation of the carbon of these com- pounds, this is an appropriate place to treat briefly of the subject of illumination from a chemical point of view, as well as that of flame, and the blowpipe, which gives a use- ful form of flame much used in the laboratory. In all kinds of illumination, except the electric light, we are dependent upon flames for the light. Whether we use illuminating gas, a lamp, or a candle, the light comes from a flame. In the first case, the gas is burned directly; in the case of the lamp, the oil is first drawn up the. wick, then converted into a gas, and this burns; while, finally, in the case of the candle, the solid material of the candle is first melted, then drawn up the wick, converted into gas, and the gas burns, forming the flame. In each case we have, then, to deal with a burning gas, and this burn- ing gas is called a flame. Illuminating Gas, Coal-gas. — Illuminating gas is some- times made from coal by heating it in closed retorts. As has already been explained, coal, particularly the softer kinds, contains compounds of carbon and hydrogen, together with some nitrogen and other elements. "When it is subjected to destructive distillation, as in the manu- facture of coal-gas, the hydrogen passes off partly in com- bination with carbon, as hydrocarbons, and partly in the 372 ILLUMINATING GAS— FLAMES. 373 free state. The nitrogen passes off as ammonia, and a large percentage of the carbon remains behind in the retort in the uncombined state as coke. The gases given off are purified, and form illuminating gas. One ton of coal yields on an average 10,000 cubic feet of gas. The value of a gas depends upon the amount of light given by the burning of a definite quantity. This is measured by comparing it with the light given by a candle burning at a certain rate. The standard candle is one made of spermaceti, which burns at the rate of 120 grains per hour; that is to say, a candle which, burning under ordi- nary conditions, loses 120 grains in one hour. The standard burner used for the gas is one through which five cubic feet of gas pass per hour. Now, to determine the illuminating power of a gas, it is passed through the standard burner at the rate mentioned, and the light which it gives is compared with the light given by the standard candle. This comparison is easily made by means of an instrument called the photometer. The illuminating power of the gas is then stated in terms of the standard candle. The statement that the illuminating power of a gas is fourteen candles, signifies that, when burning at the rate of five cubic feet per hour, its flame gives fourteen times as much light as that of the standard candle. Flames. — Ordinarily a flame is a gas which is combining with oxygen. The hydrogen flame is simply the phenome- non accompanying the act of combination of the two gases hydrogen and oxygen. Owing to the fact that we are surrounded by oxygen, we speak of hydrogen as the burn- ing gas. How would it be if we were surrounded by an atmosphere of hydrogen ? Plainly, oxygen would then be a burning gas. If we allow a jet of oxygen to escape into a vessel containing hydrogen, a flame will appear where the oxygen escapes from the jet, if a light is applied. This is an experiment which requires special precautions, and, as the principle can be illustrated as well by means of illuminating gas, this may be used instead. Just as 374 COLLEGE CHEMISTRY. illuminating gas burns in an atmosphere of oxygen, so oxygen burns in an atmosphere of illuminating gas. Kindling Temperature of Gases. — In studying the action of oxygen upon other substances, we learned that it is necessary that each of these substances should be raised to a certain temperature before it will combine with oxygen. This statement is as true of gases as of other substances. When a current of hydrogen is allowed to escape into the air, or into oxygen, no action takes place unless it is heated up to its burning temperature, when it takes fire and continues to burn, as the burning of one part of the gas heats the part which follows it, and hence it is heated to the burning temperature as fast as it escapes into the air. If the gas is cooled even very slightly below this temperature, it is extinguished. This can easily be shown by bringing down upon the flame of a Bunsen burner a piece of wire gauze. There will be no flame above the gauze, but gas will pass through unburned, and this will burn if it is lighted above the gauze. In this case, by simply passing through the thin wire gauze, the gases are cooled below their burning temperatures, and the flame does not pass through. So, also, if the gas is turned on and not lighted, and the gauze held an inch or two above the outlet, the gas will burn above the gauze if lighted above, and will not pass downward through the gauze, unless this becomes very hot. Miner's Safety-lamp — The principle illustrated in the experiments referred to in the last paragraph is utilized in the miner's safety -lamp, to which reference has already been made. One of the dangers which the coal-miner has to encounter is the occurrence in the mines of fire-damp, or methane, CH 4 , which with air forms an explosive mix- ture. The explosion can only be brought about by contact of flame with the mixture. In order to avoid the contact, the flame of the safety-lamp is surrounded by wire gauze, as shown in Fig. 69. When a lamp of this kind is brought into an explosive mixture of marsh-gas and air, the mixture passes through the wire gauze and comes in STRUCTURE OF FLAMES. 375 eontact with the flame, and a small explosion or a series of small explosions inside the gauze occurs, but the flame of the burning gas inside the wire gauze cannot pass through and raise the tem- perature of the gas outside to the burning temperature. Hence no serious explosion can take place. The flickering of the flame of the lamp, and the occurrence of small explosions inside, furnish the miner with the information that he is in a dan- gerous atmosphere. While the safety-lamp does undoubtedly afford much protection, still explosions occur. These have been shown to be caused by the presence of coal-dust in the mines, and by the com- motion of the air produced in blasting. By the aid of the coal-dust, and by sudden and violent movements of the air, it is possible for a flame surrounded by wire gauze to explode a mixture of marsh- gas and air on the other side of the gauze. Structure of Flames. — The hydrogen flame consists of a thin envelope of burning hydrogen enclosing unburned gas, and surrounded by water-vapor, which is the product of the combustion. The structure of other flames depends upon the complexity of the gases burned, and the condi- tions under which the burning takes place. In general, a flame consists of an outer envelope of gas combining with oxygen, and hence hot, and an inner part which contains unburned gas, which is comparatively cool. A part of the unburned gas is, however, quite hot, and it would combine with oxygen were it not for the fact that it is surrounded by an envelope that prevents access of air. The outer hot part of the flame is called the oxidiz- ing flame, because it presents conditions favorable to the oxidation of substances introduced into it. The inner hot part is called the reducing flame, because it consists of highly-heated substances which have the power to combine with oxygen; and hence many compounds containing 376 COLLEGE CHEMISTRY. oxygen lose it, or are reduced, when introduced into this part of the flame. The hottest part of the flame is about half-way between the bottom and the top. Here oxidation is taking place most energetically. The hottest part of the unburned gases is at the tip of the dark central part of the flame. In the flame of a Bunsen burner the two parts can be easily distinguished. The dark central part of the flame extends for some distance above the outlet of the burner. If the holes at the base of the burner are partly closed, the tip of the central part of the flame becomes luminous. This luminous tip is most efficient . ' for the purpose of reduction. The prin- jf— — 3 cipal parts of the flame are those marked in Fig. 70. The part marked I is the central cone of unburned gases; that marked c is , the luminous tip, the best part of the flame for reduction. A is the envelope of burn- ing gas. The hottest part of the flame is —i at a; that which is most efficient in causing / V \ oxidation is at d. This is further sur- I II I rounded by a non-luminous envelope con- I 11 \ sisting of the products of combustion, II carbon dioxide and water- vapor. Certain Via 70 metals placed in the upper end of the flame take up oxygen, because they are highly heated in the presence of oxygen. Certain oxides lose their oxygen when placed in the tip of the central cone, because the gases are here heated to the temperature at which they have the power to combine with oxygen. Blowpipe. — The oxidizing and reducing flames are fre- quently utilized in the laboratory. For the purpose of increasing their efficiency a blowpipe is used.- This is a tube with a convenient mouth-piece and a nozzle with a small opening through which air is blown into a flame. The blowpipe may be used with the flame of a candle, an alcohol-lamp, or a gas-lamp. It is commonly used with a gas-lamp. By regulating the current of air and slightly changing the position of the tip of the blow- CAUSES OF THE LUMINOSITY OF FLAMES. 377 pipe a good oxidizing flame or a good reducing flame can be produced. Some oxides are very easily reduced when heated in the reducing blowpipe flame. Others are not. We can frequently judge of the composition of a substance by heating in the blowpipe flame, and noticing its conduct. Some metals are easily oxidized in the oxidizing flame. Some form characteristic films, or thin layers of oxides, on the substance upon which they are heated, which is usually charcoal; and, in some cases, it is possible to detect the presence of certain substances by the color of the film of oxide. The blowpipe is therefore of much value as affording a method for the detection of the presence of certain elements in mixtures or compounds of unknown composition. The chemical principles involved in its use will be clear from what has already been said. Causes of the Luminosity of Flames. — It is evident from what we have seen that flames differ greatly in their light-giving power. The hydrogen flame, for example, though extremely hot, gives practically no light. This is also the case with the flame of the Bunsen burner; while, on the other hand, the flame of coal-gas, burning under ordinary circumstances, and that of a candle, etc., give light. To what is the difference due ? This subject has been studied very thoroughly, and it has been found that there are several causes that operate to make a flame luminous and vice versa. In the first place, if a solid sub- stance which does not burn is introduced into a non- luminous flame, a part of the heat appears as light. This is seen when a spiral of platinum wire is introduced into a hydrogen flame. It is also seen when a piece of lime is introduced into the hot non-luminous flame of the oxy- hydrogen blowpipe. A similar cause operates in ordinary gas-flames to make them luminous. Particles of unburned carbon are always present, as can be shown by putting a piece of porcelain or any solid substance into the flame, when there will be deposited on it a layer of soot, which consists mainly of finely-divided carbon. In the flame sv.ch particles are heated to incandescence, or to the tern- 378 COLLEGE CHEMISTRY. perature at which they give light. Again, it has been found that a candle gives more light at the level of the sea than it does when at the top of a high mountain, as Mount Blanc, on which the experiment was actually performed. This is partly due to a difference in the density of the gases. Naturally, the denser the gas the more active the combustion, the greater the heat, and the brighter the light. This last statement ceases to be true when the oxidation becomes sufficient to burn up all the solid particles in the flame. If gases, which in burning give light, are cooled down before they are burned, the luminosity is diminished, and, conversely, non-luminous flames may be rendered luminous by heating the gases before burning them. Gases which otherwise give lumin- ous flames give non-luminous flames when diluted to a sufficient extent with neutral gases, such as nitrogen and carbon dioxide, which neither burn nor support combus- tion. Lewes has recently shown that acetylene (which see) is always formed in flames and his work makes it appear probable that the luminosity of most, if not of all, flames is due to the formation and decomposition of this gas. Bunsen Burner. — All the statements made in regard to the causes of the luminosity of flames are based upon carefully-performed experiments. These experiments, however, cannot, for the most part, be readily repeated by the student in the laboratory in a satisfactory way. One constant reminder of the possibility of rendering a luminous flame non-luminous, and vice versa, is furnished by the burner universally used in chemical laboratories, and called, after the inventor, the Bunsen burner. The construction of this burner is easily understood. It con- sists of a base and an upper tube. The base is connected by means of a rubber tube with the gas-supply. The gas escapes from a small opening in the base, and passes upward through the tube. At the lower part of the tube there are two holes, which may be opened or closed by turning a ring with two corresponding holes in it. When the gas is turned on, it is lighted at r 1 ' *. the tube. CYANOGEN. 379 Air is at the same time drawn through the holes at the base. The result is that the flame is practically non- luminous. If the ring at the base is turned so that the air-holes are closed, the flame becomes luminous. The advantage of the non-luminous flame for laboratory use consists in the fact that it does not deposit soot, and, at the same time, it is hot. The non-luminosity of the flame of the Bunsen burner appears to be due to several causes: (1) Dilution of the gases by means of the nitrogen of the air; (2) Cooling of the gases by the entrance of the air; (3) Burning of the solid particles by the aid of the oxygen of the air admitted to the interior of the flame. Compounds of Carbon with Nitrogen and with Sulphur. Cyanogen, C 2 N 2 . — Carbon does not combine with nitro- gen under ordinary circumstances. If, however, these elements are brought together at very high temperatures in the presence of metals, they combine to form com- pounds known as cyanides. Thus, when nitrogen is passed over a highly-heated mixture of carbon and potassium carbonate, potassium cyanide, KCN, is formed. Carbon containing nitrogen, as animal charcoal, when ignited with potassium carbonate, reduces the carbonate, forming- potassium, in presence of which carbon and nitrogen com- bine, forming potassium cyanide. When refuse animal substances, such as blood, horns, claws, hair, etc., are heated together with potassium carbonate and iron, a sub- stance known as potassium ferrocyanide, or yellow prussiate of potash, K 4 Fe(CN) 6 + 3H 2 0, is formed. When this is simply heated it is decomposed, yielding potassium cyanide: K 4 Pe(CN) 6 = 4KCN" + PeC a + N 2 . It is an easy matter to make the mercury salt, Hg(CN) 2 , 380 COLLEGE CHEMISTRY. from the potassium salt. By heating mercuric cyanide it breaks up, yielding metallic mercury and cyanogen gas : Hg(CN) 2 = Hg + C 2 N 2 , just as mercuric oxide yields mercury and oxygen when heated : HgO = Hg+0. Cyanogen (from xvaros, Hue) owes its name to the fact that several of its compounds have a blue color. It is a colorless gas, which is easily soluble in water and alcohol, and is extremely poisonous. It burns with a purple-colored flame. Hydrocyanic Acid, Prussic Acid, HON. — This acid, which is commonly called prussic acid, occurs in nature in amygdalin, in combination With other substances, in bitter almonds, the leaves of the cherry, laurel, etc. It is prepared by decomposing metallic cyanides with hydro- chloric acid. It is volatile and passes over. The action is represented thus : KCN + HC1 = KC1 + HON. Hydrocyanic acid is a volatile liquid, boiling at 26.1°, and solidifying at — 14°. It has a characteristic odor suggestive of bitter almonds. It is extremely poisonous. It dissolves in water in all proportions. Such a solution is known as prussic acid. Pure hydrocyanic acid is very unstable. By standing, a brown substance is deposited from its solution. Cyanic Acid, HONO. — By gentle oxidation of a cyanide it is converted into a cyanate. Thus, by melting together potassium cyanide and lead oxide, potassium cyanate is formed : KCN + PbO = KOTO + Pb. Cyanic acid is a volatile; acrid, unstable liquid. It CARBON DISULPHIDE. 381 breaks down at once into carbon dioxide and ammonia in presence of water : CONH + H 2 = NH 3 + C0 2 . The potassium salt is easily soluble in water, but is decomposed by it, yielding ammonia and acid potassium carbonate : CONK + 2H 2 = KHCO s + NH 5 . These decompositions of cyanic acid and the cyanates further exemplify the tendency of cyanogen compounds to undergo decomposition in presence of water. Carbon Disulphide, CS 2 . — Just as carbon combines directly with oxygen to form the dioxide, so it combines directly with sulphur to form the disulphide; but there is a great difference in the ease with which carbon combines with the two elements. In order to effect combination with sulphur a very high temperature is necessary. The compound is prepared on the large scale by heating char- coal to a high temperature in an upright cast-iron cylinder, and adding sulphur in such a way that it enters the bottom of the cylinder. The product is passed through a series of tubes arranged so as to secure condensation. Carbon disulphide is a clear liquid which has a high refractive power. It boils at 46.2°. When pure it has a pleasant odor, but if kept for a time, particularly if water is present in the vessel, it undergoes slight decomposition, and products of extremely disagreeable odor are formed. It can generally be freed from these by shaking it with a little mercury and then redistilling. It burns readily, forming carbon dioxide and sulphur dioxide: CS 2 + 30 2 = C0 2 + 2S0 2 . Carbon disulphide is only slightly soluble in water, and is decomposed by it only very slowly. The disulphide is 382 COLLEGE CHEMISTRY. an excellent solvent for many substances ■which are not soluble in water, as, for example, fats, resins, iodine, and one of the modifications of sulphur and of phosphorus. The solution of iodine in it has a beautiful violet color; and when a water solution containing a little free iodine is shaken with carbon disulphide the latter acquires a violet color and separates below the water. Carbon disulphide finds extensive application as a solvent, and it is also used for the purpose of destroying phylloxera, the insect which is so destructive to grape-vines, particularly in the wine districts of France. Sulphocyanic Acid, HCNS. — Just as the cyanides take up oxygen and are converted into cyanates, so also they take up sulphur and are converted into sulphocyanates : KON + S = KCNS. Constitution of Cyanogen and its Simpler Compounds. — The compounds of cyanogen show, in general, a remark- able similarity to the compounds of the chlorine group. The hydrogen compound is a monobasic acid and forms a series of salts, the cyanides, which in general are analogous to the chlorides. Comparing the cyanides with the chlorides it is clear that in the former the group (CN), or the cyanogen group, plays the same part that the atom chlorine plays in the chlorides : H(CN) HC1 K(CN) KC1 Hg(CN) 2 HgCl a . So, also, cyanic acid and hypochlorous acid are analogous : HO(CN) HOC1. This relation suggests that which is observed between the ammonium compounds and those of potassium and sodium. The cyanogen group is evidently univalent, as it combines with one atom of hydrogen, one of potassium, etc. EXPERIMENTS : COAL-GAS, ETC. 383 EXPERIMENTS. COAL-GAS. Experiment 186. — Heat some bituminous coal in a retort and collect over water the gases given off. Are these gases com- bustible ? Oxygen burns in an Atmosphere of a Combustible Gas. Experiment 187.— Break off the neck of a good-sized retort ; fit a perforated cork to the small end ; pass a piece of glass tube through the cork, and connect by means of rubber hose with an outlet for coal-gas. Fix the apparatus in position, as shown in Fig. 71. Turn the gas on, and when the air is driven out of the Fig. 71. retort-heck, light the gas. The neck is now filled with illn- minating gas, and the gas is burning at the mouth of the vessel. If now a platinum jet from which oxygen is issuing is passed up into the gas the oxygen will take fire, and a flame will appear where the oxygen escapes from the ]et. The oxygen " burns " in the atmosphere of coal-gas. Kindling Temperature of Gases. Experiment 188. — Light a Bunsen burner. Bring down upon the flame a piece of brass or iron wire-gauze. There is no flame above the gauze. That the gas passes through unburned can be shown by applying a light just above the outlet of the burner 3«4 COLLEGE CHEMISTRY. and above the gauze. The gas will take fire and burn. By simply passing through the thin wire-gauze, then, the gas is cooled down below its burning temperature, and does not burn unless it is heated up again. Turn on a Bunsen burner. Do not light the gas. Hold apiece of wire- gauze about one and a half to two inches above the outlet. Apply a lighted match above the gauze, when the gas will burn above the gauze, but not below it. Here again the heat necessary to raise the temperature of the gas to the burning temperature cannot be communicated through the gauze. If in either of the above-described experiments the gauze is held in position for a time, it will probably become so highly heated that the gas on the side where there is no flame will be raised to the burning tempera- ture. The instant that point is reached the flame becomes continuous. 0- The Blowpipe and its Uses. Fig. 73. The blowpipe used in chemical laboratories is con- structed as shown in Fig. 72. "When used with the Bunsen burner it is best to slip into the Darner a brass tube ending above in a narrow slit-like opening, as shown in Fig. 73. The tube re- ferred too, marked a in the figure, reaches to the bottom of the burner, and thus cuts off the supply of air which usually enters the holes at the base. The gas is now lighted, and the current so regulated that there is a small flame about 1£ to 2 inches long. " The tip of the blowpipe is placed on the Fig. 73. Fig. 74. (Slit of the burner in the flame, as shown in Fig. 74. By blowing THE BLOWPIPE AND ITS USES. 3 8 S regularly and not violently through the pipe the flame is forced down in the same direction as the end-piece of the blowpipe, and the slant of the burner-slit. Under proper conditions the flame separates sharply into a central blue part and an outer part of another color. The direction and lines of division of the flame are indicated in Fig. 74. The outer part of the flame marked o is the oxidizing flame ; the part marked r is the reducing flame. Experiment 189. — Select a piece of charcoal about 4 inches long by 1 inch wide and 1 inch thick, with one surface plane.* Near the end of the plane surface make a cavity by pressing the edge of a small thin coin against it, and turning it completely round a few times. Mix together equal small quantities of dry sodium carbonate and lead oxide. Put a little of the mixture in the cavity in the charcoal, and heat it in the reducing flame produced by the blowpipe. In a short time globules of metallic lead will be seen in the molten mass. After cooling, scrape the solidified substance out of the cavity in the charcoal. Put it in a small mortar, treat it with a little water, and, after breaking it up and, allowing as much as possible to dissolve, pick out the metallic beads. Is it malleable or brittle ? Is metallic lead mal- leable or brittle ? Is it dissolved by hydrochloric acid ? Is lead soluble in hydrochloric acid ? Is it soluble in nitric acid ? Is lead soluble in nitric acid ? The action of the acids can be tried by putting a bead on a small dry watch-glass and adding a few drops of the acid. Does the substance act like lead ? What has become of the oxygen with which the lead was combined in the oxide ? Is there any special advantage in having a support of charcoal for this experiment ? Experiment 190. — Heat a small piece of metallic lead on char- coal in the oxidizing blowpipe flame. Notice the formation of the oxide, which forms a coating or film on the charcoal in the neighborhood of the metal. Is there any analogy between this process and the burning of hydrogen ? In what does the analogy consist ? What differences are there between the two processes ? A Experiment 191. — Repeat the experiments with arsenic, anti- mony, and bismuth. Notice the colors of the films formed on i the charcoal. \ Experiment 192. — Melt into a bit of glass tubing a piece of platinum wire 8 to 10 mm. (3 to 4 inches) long, and bend the end so as to form a small loop, as shown in Eig. 75. Heat the loop in the flame of a Bunsen burner, and then dip it into some sodium- * Pieces of charcoal prepared for blowpipe work can be bought from dealers in chemical apparatus at small cost. 386 COLLEGE CHEMISTRY. ammonium phosphate (microcosmic salt). Heat in the oxidizing flame of the blowpipe until a clear glass bead is formed in the loop. What changes have taken place? and what is the clear T> Fig. 75. glass ? Bring a minute particle of a manganese compound in contact with the bead, and heat again. What change takes place ? Try the same experiment, using successively a cobalt compound, a copper compound, and an iron compound. Now, instead of using microcosmic salt, use borax. Explain the changes in all the above-described experiments. Cyanogen. Experiment 193. — Make potassium cyanide by heating potas- sium ferrocyanide in an iron crucible. Experiment 194. — Make cyanogen by heating mercuric cy- anide. Cyanogen is poisonous. Burn some of the gas. Experiment 195. — Make potassium cyanate from some of the cyanide obtained in Experiment 193. This is done by melting it in an iron crucible, and, while the mass is liquid, adding about four times its weight of red lead, stirring- during the operation. After this the crucible should again be put in the furnace for a little while, the metallic lead allowed to settle, and the contents poured out on a smooth stone. Break this up, and extract the cyanate with alcohol. CHAPTER XXII. ELEMENTS OF FAMILY IV, GROUP A: SILICON.— TITANIUM.— ZIRCONIUM.— CERIUM.— THORIUM. General. — While in some respects the elements of this group resemble carbon and bear to it relations similar to those which the members of the chlorine group bear to fluorine, the members of the sulphur group to oxygen, and the members of the phosphorus group to nitrogen, yet between them and carbon there are some remarkable differences. All the members of the group except titanium combine with hydrogen. The compounds formed have the formulas SiH 4 , ZrH 2 , CeH 2 , and ThH 2 . All the elements of the group form oxygen compounds analogous to carbon dioxide. They are : Si0 2 , Ti0 2 , Zr0 2 , CeO a , ThO, The first three are acidic, and form salts which in com- position are analogous to the carbonates. These are the silicates, titanates, and zirconates of the general formulas M 2 Si0 3 , M 2 Ti0 3 , M 2 Zr0 3 . Cerium and thorium oxides are basic. These facts sug- gest the relations between the members of the phosphorus group. The oxides of the last two members, antimony and bismuth, are basic, although the oxide of antimony is also acidic in its conduct towards the stronger bases. 387 388 COLLEGE CHEMISTRY. The compounds of silicon are very abundant in nature; those of the other members of the group are rare. Silicon, Si (At. Wt. 28.4). Occurrence. — We have already seen what an exceedingly important part carbon plays in animate nature. It is in- teresting to note that silicon, which in some respects from a chemical point of view resembles carbon, is one of the most important constituents of the mineral or inorganic parts of the earth. It occurs chiefly in the form of the dioxide, SiO a , commonly called silica, or silicon dioxide; and in combination with oxygen and several of the common metallic elements, particularly with sodium, potassium, aluminium, and calcium, in the form of the silicates. Next to oxygen, silicon is the most abundant element in the earth. There are extensive mountain-ranges consist- ing almost entirely of the dioxide, Si0 2 , in the form known as quartz or quartzite. Other ranges are made up of silicates, which are compounds formed by the combina- tion of silicon dioxide and bases. The clay of the valleys and river-beds also contains silicon in large quantity, while the sand found so abundantly on the deserts and at the seashore is largely silicon dioxide. Preparation. — Silicon does not occur in nature in the free state. The oxide, SiO s , which is most abundant in the form of sand, is decomposed by heating it with potas-: sium or magnesium, and silicon is thus set free. When magnesium is used the action is violent, and besides the silicon a compound of silicon and magnesium is forced. Silicon has also been made by heating the oxide and carbon in the electric furnace, and by decomposing the chloride with potassium: Sid. + 4K = Si + 4BC1. The best way to make it is byheating together potassium fluosilicate, K 2 SiF 6 , sodium, mMt zinc : K.Sil', + 4Na = 4NaF + 2KE + Si. SILICON. 3 8 9 At the same time the zinc melts and the silicon which separates dissolves in the molten zinc. On cooling, it is deposited from the solution in beautiful needle-shaped crystals, around which the zinc solidifies at a lower tem- perature. By treating the mass with hydrochloric acid the zinc is dissolved and the crystals of silicon are left behind. When obtained by reduction of the oxide or the chloride by means of potassium, it is a brown amorphous powder. If made by decomposition of potassium fluosili- cate by aluminium, it is deposited from the molten aluminium in crystals somewhat resembling graphite. Just as there are three forms of carbon, the amorphous, graphite, and diamond, so there are three corresponding forms of silicon, the amorphous brown powder, the graphi- toidal, and the needles. The amorphous variety is con- verted into crystallized silicon by continued heating at a high temperature. Amorphous silicon acts upon hydrofluoric acid, forming silicon tetrachloride, SiF 4 , and setting hydrogen free: Si + 4HP = SiF, + 2H a . In this reaction it exhibits one of the properties of a base- forming element. Towards other acids, however, it is indifferent. It is not acted upon by sulphuric acid, nor by nitric acid, nor aqua regia. It dissolves, however, in potassium hydroxide, forming potassium silicate, in this case acting like an acid-forming element : Si + 2KOH + H 3 = K 2 Si0 3 + 2H 3 . This form of silicon also burns in the air, forming the dioxide. Crystallized silicon, on the other hand, does not burn in oxygen at the highest temperatures. It, however, reduces carbon dioxide and decomposes carbonates at a high temperature. It is also oxidized by a melting mix- ture of potassium nitrate and the hydroxide or carbonate. It combines with nitrogen at a high temperature. 39° COLLEGE CHEMISTRY. Both the graphitoidal and needle-formed crystals of silicon consist of regular octahedrons. Both forms have a blacMsh-gray color and a metallic lustre. Silicon Hydride, SiH 4 . — This gas is obtained mixed with hydrogen when a compound of magnesium and silicon is treated with hydrochloric acid : Mg 2 Si + 4HC1 = SiH 4 + 2MgCl 2 . Thus made, it takes fire when it comes in contact with the air, and the act is accompanied by explosion. The products of its combustion are silicon dioxide and water. When pure it forms a colorless gas which does not take fire spontaneously in the air at the ordinary temperature. If it is diluted with hydrogen, or if it is heated, it does take fire. When burned in a cylinder or narrow tube, so that free access of air is not possible, amorphous silicon is deposited upon the walls of the vessel. Titanium, Ti (At. Wt. 48.1). — Titanium occurs in nature as titanium dioxide, Ti0 2 , in three distinct forms, known as rutile, brookite, and anatase; in combination with iron, as titaniferous iron which contains ferrous titanate, PeTi0 3 ; and in a number of iron ores and rare minerals. Zirconium, Zr (At. Wt. 90.7). — The principal form in which zirconium occurs in nature is as zircon, which is a silicate of the formula ZrSi0 4 , derived from normal silicic acid, Si(OH) 4 , by the substitution of a quadrivalent atom of zirconium for the four hydrogen atoms. Thorium, Th (At. Wt. 232.5).— This element occurs principally in the mineral thorite, which is essentially a silicate of thorium, ThSi0 4 , analogous to zircon. Cerium so much resembles the two elements lanthanum and didymium that, although it falls in the same group as silicon and resembles the elements of this group in some respects, it seems advisable to postpone its study until lanthanum and didymium are taken up. compounds of the silicon group. 391 Compounds of the Elements of the Silicon G-koup with those of the Chlorine Group. Silicon Tetrachloride, Si01 4 . — This compound is formed when silicon is heated in. a current of chlorine, and by passing a current of dry chlorine over a heated mixture of silicon dioxide and carbon. Under these latter circum- stances the following reaction takes place : Si0 2 + 2G + 2C1 2 = SiCl 4 + 200. The tetrachloride is a colorless liquid. It is decomposed by water, forming silicic acid and hydrochloric acid. The reaction probably takes place as represented in the follow- ing equation: SiCl 4 + 4H 2 = Si(OH) 4 + 4HC1. The normal acid thus formed breaks down readily, how- ever, forming the ordinary acid of the formula SiO(OH) 2 or H 2 SiO s , corresponding to carbonic acid, H 2 C0 3 . Silicon Hexachloride, Si 2 Cl 6 , is formed -when silicon tetrachloride is heated with silicon : 3SiCl 4 + Si = 2Si 2 Cl 6 . When heated to a sufficiently high temperature it is decomposed, yielding silicon and the tetrachloride : 2Si 2 Cl 6 = 3SiCl 4 + Si. Water decomposes it, forming the corresponding hydroxyl derivative, which loses water and forms the acid Si 2 3 (OH) 2 : Si 2 Cl 6 + 6H 2 = Si 2 (OH) 6 + 6HC1; Si 2 (OH) 6 = Si 2 2 (OH) 2 + 2H 2 0. The product is a disilicic acid, in some respects analogous to disulphuric acid. 39 2 COLLEGE CHEMISTRY. Silicon Tetrafluoride, SiF 4 . — This is one of the most interesting of the compounds that silicon forms with the members of Family VII. It is made by treating silicon dioxide with hydrofluoric acid. This action is secured by treating a mixture of silicon dioxide (sand) and calcium fluoride (fluor-spar) with concentrated sulphuric acid, when two reactions take place: CaF 2 + H 2 S0 4 = CaS0 4 + 2HF; SiO a + 4HF = 2H 2 + SiF 4 . The tetrafluoride escapes as a colorless gas, that forms thick clouds in moist air on account of the action of water upon it. Water decomposes the tetrafluoride, as it does the tetra- chloride. The first action probably consists in the forma- tion of normal silicic acid and hydrofluoric acid, the normal acid then breaking down by loss of water and yielding the ordinary form of silicic acid : SiF 4 + 4H 2 = Si(OH) 4 + 4HF; Si(OH) 4 = SiO(OH) 2 + H 2 0. The silicic acid thus formed separates as a gelatinous mass. At the same time the hydrofluoric acid acts upon some of the silicon tetrafluoride, forming the compound jiuosilicic acid, which has the formula H 2 SiF 6 : SiF 4 + 2HF = H 2 SiF 6 . The complete action may be represented in one equation, as follows: 3SiF 4 + 3H 2 = H 2 SiO s + 2H 2 SiF 6 . The fluosilicic acid remains in solution in the water, and by treating this solution with carbonates or hydroxides of the metallic elements the salts known as the fluosilicates are obtained. The solution of the acid can be concentrated to a certain extent in a platinum -vessel, but it breaks down COMPOUNDS OF THE SILICON GROUP. 393 into silicon tetrafluoride and hydrofluoric acid when it becomes concentrated. If more potassium hydroxide than is required to neutralize the acid is added to the solution, decomposition ensues, with formation of silicic acid : H 2 SiF e + 6KOH = 6KF + H 2 SiO s + 3H,0. By water alone, however, the acid is not decomposed, and the salts are fairly stable. When heated, the salts give off silicon tetrafluoride, and fluorides are left behind : K 2 SiF 8 = 2KF + SiF,. Compounds of the Members of the Silicon Group with Oxygen, and with Oxygen and Hydkogen. Silicon Dioxide, SiO z . — This compound occurs very abundantly in nature in many different forms, both crystallized and amorphous. Quartz is a form of crystal- lized silicon dioxide which is found very widely distributed. It crystallizes in the hexagonal system in prisms and pyramids, the crystals sometimes attaining great size and beauty. Another form of the crsytallized compound is that known as tridymite. Like quartz it crystallizes in the hexagonal system, but the characteristic forms are not the same as those of quartz. Further, it nearly always occurs in triplet crystals. The finer crystals of quartz are generally called rock-crystal; the crystalline variety in which the crystals are not well developed is called quartz- ite. The amorphous varieties of silicon dioxide frequently contain water in combination, or, rather, they are hy- droxides of silicon. Examples of these forms are opal, agate, amethyst, carnelian, flint, sand, chalcedony, Some of these are colored by small quantities of other substances contained in them. Carnelian ower its color to a com- pound of iron, probably ferric oxide; flint contains small quantities of organic matter. The specific gravity of the 394 COLLEGE CHEMISTRY. crystallized varieties is higher than that of the amorphous varieties, and there are also some chemical differences between them. Properties. — Silicon dioxide is insoluble in water and in most acids. It dissolves, however, in hydrofluoric acid, forming the tetrafluoride. It requires the temperature produced by the oxyhydrogen blowpipe to melt it. The amorphous varieties are more easily acted upon by other substances than the crystallized. Thus, hydrofluoric acid acts much more readily upon them. When the amorphous compound is boiled with solutions of potassium or sodium hydroxide, or of the carbonates of these metals, it dis- solves, forming the corresponding silicate: K 2 CO s + Si0 2 = K a SiO s + C0 2 ; 2KOH + Si0 2 = K a SiO s + H 2 0. The crystallized varieties are not dissolved in this way. All forms of the dioxide act upon melting hydroxides or carbonates of potassium or sodium, and form the corre- sponding silicates. ff Uses. — Plants take up silicon dioxide from the soil, and this being deposited in various parts of their tissues, gives them the necessary firmness. Straw, for example, is rich in silicon dioxide. Horsetail, a plant of the genus Equisetum, is so rich in finely-divided silicon dioxide that it is used for polishing. There are great natural deposits of finely-divided silicon dioxide known as infusorial earth. This consists of the remains of diatoms. And finally sili- con dioxide is found in the hair, in feathers, and in egg albumen. Silicon dioxide finds extensive application in the manufacture of mortar, glass, and porcelain. Ordinary glass, as we shall see, is a silicate of calcium and potassium or sodium, which is made by melting together sand and the carbonates of the metals mentioned. Silicic Acid. — There are many varieties of silicic acid, all of which can, however, be referred to the normal acid, SILICIC ACID. 395 Si(0H) 4 . This normal acid is contained in the gelatinous precipitate which is formed when silicon tetrachloride or tetrafluoride is decomposed by water : SiCl, + 4H 3 = Si(0H) 4 + 4HGL . \^\ C\\e sleV The normal acid readily loses water and forms the acid of the formula OSi(OH) 2 or H 2 Si0 3 . Prom the latter most of the ordinary silicates are derived. It cannot be isolated, for when filtered off and exposed to the air it loses more water, and when heated to a sufficiently high temperature it is converted into silicon dioxide. When potassium or sodium silicate in solution is treated with hydrochloric acid, most of the silicic acid separates in the form of a gelatinous mass if the solution is concen- trated. If, however, the solution is dilute, a considerable part of the acid remains in solution. Further, if a con- centrated solution of the silicate of potassium or sodium is poured quickly into hydrochloric acid, or if the acid is poured quickly into the solution of the silicate, the silicic acid remains in solution. If, however, the solutions are brought together drop by drop the silicic acid separates. Prom these solutions of silicic acid ammonia or ammonium carbonate throws down the acid. A solution of pure silicic acid can be obtained by means of dialysis. It has been found that solutions of different substances pass with different degrees of ease through porous membranes, just as gases differ as regards the ease with which they pass through porous diaphragms. This fact concerning gases was referred to in connection with hydrogen. Now, while some solutions pass readily through parchment paper, others pass through with difficulty, and some do not pass through at all. A dialyser, or an apparatus used in dialysis, may be made by tying a piece of parchment paper over the mouth of a ring-formed glass or rubber vessel, and placing this in another shallow vessel. Pure water is put in the outer vessel, and the 396 COLLEGE CHEMISTRY. solution for dialysis in the inner one. The arrangement is illustrated in Pig. 76. In the figure aa is the hoop of gutta-percha, and b is the parchment paper. "When now the solution containing hydrochloric acid, sodium chloride, and silicic acid is put in the dialyser, the hydrochloric acid and sodium chloride Fig. 76. pass readily through the membrane, while the silicic acid is left- behind, and in the course of a few days, if the water in the outer vessel is renewed, the solution of silicic acid in the inner vessel will be found to be free from the other substances. This solution can be evaporated to some extent by boiling, but when a certain concentration is reached the acid separates. In a vacuum such a solution can be evaporated further without the formation of a deposit. Finally, there is left a transparent mass which has approximately the composition represented by the formula H 2 SiO s . The dialysed solution of silicic acid is coagulated by a very dilute solution of sodium or potassium carbonate, and by carbon dioxide itself. When the solutions containing silicic acid are evaporated to complete dryness the acid is converted into silicon dioxide and insoluble hydrates. This residue is called insoluble silicic acid. When this is treated with hydro- chloric acid and water it remains undissolved, and if filtered off and ignited it leaves a residue of silicon dioxide. To sum up, then: Whenever silicic acid is formed in a solution it is a more or less complex derivative of normal P0LYSILIC1C AND TRISILtCtC ACIDS. 397 silicic acid, and is somewhat soluble in water, but by the processes just described the soluble acid is converted into insoluble silicon dioxide, as explained. Polysilicic Acids. — Silicic acid is remarkable for the great number of derivatives which it yields. Most of these bear to the normal acid relations similar to those which the various forms of phosphoric acid bear to normal phos- phoric acid. It has already been stated that salts of the acid H 2 Si0 3 are more common than those of the normal acid. Among the salts of the normal acid are zircon, ZrSi0 4 , and thorite, ThSi0 4 . The ordinary silicates of potassium and sodium are derived from the acid H 2 SiO } ; so also are wollastonite, CaSi0 3 , and enstatite, MgSi0 3 . Disilicic Acid is derived from ordinary silicic acid by loss of one molecule of water from two molecules of the acid : _ .OH OH 0S < 20Si 39 8 COLLEGE CHEMISTRY. the first acid is expressed by formula I, and of the second by formula II, below given : Si ) (OH), Sij°, H Si {if Si |(° H )* SijS SijS S M(OH) 3 Si{g (0 H) s Si i 8 3 A1 i. ii. in. Orthoclase or ordinary feldspar is the aluminium-potas- sium- salt of the second form of trisilicic acid, in -which one atom of hydrogen is replaced by potassium, and three by an aluminium atom, as shown in formula III above. Titanium Dioxide, Ti0 2 . — As has been stated, this is one of the principal forms in which titanium is found in nature. There are three natural crystallized varieties — rutile, brookite, and anatase. Silicides are compounds of silicon with other elements, as, for example, with carbon. These two elements com- bine, forming an interesting compound, carbon silicide, CSi, which is manufactured on the large scale and known in the market as carborundum. This is made by heating a mixture of quartz sand, coke, and common salt, or sodium chloride, in the electric furnace to 3500°, when the reaction represented below takes place : SiO s + 20 = CSi -t- 2CO. The product is in the main crystallized, the crystals being bluish or yellowish-green. They have the specific gravity 3. 22 to 3. 12. The silicide is said to be colorless when perfectly pure. It scratches ruby and chrome-steel, and on account of its hardness it is much prized as a polishing agent, being used to a considerable extent in place of emery. Pure carbon silicide is insoluble in nearly all ordinary solvents, including hydrochloric, nitric, sul- phuric, and hydrofluoric acids. It is, however, decom- posed by fusing caustic alkalies or their carbonates EXPERIMENTS WITH SILICON. 399 Many other silicides have been made, several of which are well-crystallized compounds. Family IV, Group B. • Allied to the members of the silicon group, yet differing from them in some important particulars, are the three elements germanium, tin, and lead. Of these the first two are more acidic in character than the last. They combine with chlorine in two proportions, forming the chlorides GeCl 2 , SnCl 2 , PbCl 2 , GeCl 4 , SnCl 4 , PbCl 4 . With oxygen they unite, forming the compounds Ge0 2 , Sn0 2 , and Pb0 2 . Stannic oxide, Sn0 2 , and lead peroxide, Pb0 2 , form salts with bases, and these have the composi- tion represented by the general formulas M 2 Sn0 3 and M,PbO s , and are therefore analogous to the silicates and titanates. On the other hand, further, salts are known which are derived from the oxide PbO. These have the general formula M 2 PbO a , and are to be regarded as salts of an acid, Pb(OH) a . These salts are not stable, and are not easily obtained. Most of the derivatives of lead are those in which it plays the part of a base-forming element. It will therefore be better to postpone its study until it is taken up under the general head of the base-forming tjlements. Notwithstanding, further, the marked analogy between some of the compounds of tin and those of the members of the silicon group, it appears on the whole advisable to treat of this element in company with lead, which it also resembles in many respects. f yJ^ EXPERIMENTS. Silicon. Experiment 196. — Prepare sodium fluosilicate as directed in the next experiment. Mix 3 parts of the dry salt with 1 part of sodium cut in pieces. Throw this mixture all at once into a Hes- sian crucible heated to bright red heat in a furnace. Add imme- diately 9 parts granulated zinc, and a layer of sodium chloride previously heated to drive off water. The crucible is then cov- 400 COLLEGE CHEMISTRY. ered, and the fire allowed to burn down. After cooling, the regulus of zinc containing the silicon is separated from the slag, washed with water, and treated with hydrochloric acid. The zinc dissolves and leaves the silicon. This is again washed with water, and then heated with nitric acid, and washed with water, when crystals of silicon, sometimes of great beauty, are obtained. Try the effect of heating a little of the silicon in the air. Try the action of acids and of alkalies upon it. Silicon Teteaeluoride and Fluosilicic Acid. Experiment 197— Arrange an apparatus as shown in Fig. 77. A is a bottle of about 2 litres capacity, such as are commonly used for transporting acids. This is about two-thirds filled with S^ Fig. 77. alternating layers of sand and powdered fluor-spar, moistened with concentrated sulphuric acid. The bottle is put in the deep sand-bath B, and connected by means of a wide glass tube with the funnel 0, which dips just below the surface of the water in the large evaporating dish D. The sand-bath is now gently heated, when silicon tetrafluoride passes over. Coming in con- tact with water, it is decomposed, silicic acid being deposited and fluosilicic acid passing into solution. In order to prevent clog- ging, the gelatinous silicic acid is from time to time removed from the mouth of the funnel by means of a bent glass rod. After the action is complete, filter the solution. Take out one- EXPERIMENTS WITH SILICIC ACID. 4°* quarter, and to the rest slowly add a solution of sodium carbon- ate until the whole just begins to show an alkaline reaction ; now add the other quarter of the acid, and filter. Explain all the re- actions. Heat a little of the dried salt in a covered platinum crucible. What change takes place ? What evidence have you that the change has taken place ? To a little of the salt in water add a solution of potassium hydroxide. What change takes place ? Dry the silicic acid formed in the first part of the experiment by decomposition of the silicon tetrafluoride. Silicic Acid. Experiment 198. — Boil some of the silicic acid obtained in the last experiment in a solution of sodium hydroxide. Treat some of the solution with hydrochloric acid ; with ammonium chloride. Experiment 199. — Fuse a mixture of equal weights potassium and sodium carbonates in a platinum crucible in the flame of the blast-lamp and slowly add fine sand to the molten mass. Con- tinue the heating until no more sand is dissolved. Pour the molten mass out on a stone, and when cooled break it up and treat it with water. Experiment 200. — Treat a little of the solution containing sodium and potassium silicates, prepared in the last experiment, with a little sulphuric or hydrochloric acid. A gelatinous sub- stance will be precipitated. This is silicic acid. Some of the acid remains in solution. By evaporating the solution to dryness and heating for a time on the water-bath, all the silicic acid is rendered insoluble. CHAPTEK XXIII. CHEMICAL ACTION. S^' Retrospective. — "We have been studying the principal elements of four families and the compounds which they form with one another. No matter how simple or how complex the chemical changes studied were, certain funda- mental laws governing all cases of chemical action were found to hold good. These laws have been discussed, but it will be well to recall them here before taking up others which are intimately connected with them. The first great law of chemical change is I. The law of conservation of mass. According to this the amount of matter is not changed by a chemical act. The second law is II. The law of definite proportions. According to this, the composition of any compound v always the same. The third law is III. The law of multiple proportions. According to this, the different masses of any element that combine with a fixed mass of another or others bear simple relations to one another. To account for the laws of definite and multiple propor- tions the Atomic Theory has been proposed. According to this, each element is made up of particles of definite weight, which are chemically indivisible, and 402 CHEMICAL ACTION. 4°3 chemical action consists in union or separation of these particles. These hypothetical particles are called atoms. The elements must combine in the proportion of their atomic weights or of simple multiples of these, if the atomic theory is true. Further study showed that it is necessary to assume the existence of larger particles than the atoms, viz., the mole- cules. According to the theory of molecules, every chem- ical compound and element is made up of molecules, which are the smallest particles having the same general proper- ties as the mass. These molecules are made up of atoms which, in the case of compounds, are of different kinds, and in the case of elements, of the same kind. In the case of a few elements the atom appears to be identical with the molecule. From the study of gases the conclusion is reached that in equal volumes of all gases under standard conditions there is always the same number of molecules (Avogadro's hypothesis). This furnishes a means of determining the relative weights of molecules of gaseous substances ; and from these molecular weights it is possible to draw conclusions in regard to the atomic weights of those elements which enter into the composition of the compounds thus studied. The formulas of chemical compounds are intended to be molecular formulas. They are intended to tell of what atoms and of how many atoms the molecules represented are made up. The method of determining molecular weights based upon Avogadro's hypothesis is applicable only to gaseous substances, or to such as can be converted into gas without undergoing decomposition. While many of the com- pounds with which we have had to deal are of this character, many of them ..are not, and in regard to the molecular weights of these we must be in doubt unless some other method applicable to liquids and solids is available. So, too, the atomic weights of those elements which enter into the composition of gaseous compounds can be deduced from the molecular weights, but plainly 4°4 COLLEGE CHEMISTRY. those which do not enter into the composition of such compounds demand some other method. For determining the atomic weights of such elements an excellent method is based upon the study of specific heats; while for the determination of the molecular weights of solid substances which can be dissolved without decomposition a method has been worked out that is based upon the extent to which the compound raises the boiling-point or lowers the freez- ing-point of its solution. Both these methods will be briefly described in this chapter. Next, it is found that there is a limit to the law of multiple proportions. While, according to this law, the masses of any element which unite with a given mass of another element bear simple relations to one another, the law is silent as to how many kinds of compounds are possible between any two elements. A careful examina- tion of the composition of the compounds of the elements shows, however, that there is a limit to the number of atoms of one element that can combine with one atom of another element. This limit is determined by what is called the valence of the elements. Observations on the composition of compounds have led to the hypothesis of the linking of atoms — the linking taking place according to the laws of valence. The arrangement of the atoms in a molecule is, acording to this, the constitution of a com- pound. Valence, as we have seen, is not a constant property of the atoms. Towards oxygen the elements which we have thus far studied have the highest valence; towards hy- drogen the lowest; and, in general, towards the members of the chlorine group they exhibit an intermediate valence. The valence towards hydrogen is in most cases constant, while the valence towards oxygen and towards the mem- bers of the chlorine group varies, in some cases between comparatively wide limits, as between 1 and 7 in the chlorine group, and between 2 and 6 in the sulphur group. Further, the variations in the valence of an element generally take place from odd to odd or from even to even. CLASSIFICATION OF REACTIONS. 405 In the case of chlorine it appears to vary from 1-3 to 5-7; in that of sulphur, from 2-4. to 6; in that of phosphorus, from 3 to 5. A comparison of the atomic weights finally led to the discovery that the properties of the elements are a periodic . function of these weights. This is the periodic law of chemistry. This makes a systematic classification jf the elements according to their atomic weights and their properties possible, and is full of suggestion as to the relations which the forms of matter we call elements bear to one another. Classification of Reactions of the Elements and Com- pounds Studied. — While there is undoubtedly something confusing in the number of the compounds and their reactions which we have been studying, still, when these are interpreted in the light of the atomic theory, of the law of valence, and of the periodic law, the study is much simplified, and those things which seem to have little or no connection are found to form parts of a general system. In studying chemistry, one of the first things to be done is to learn how elements and compounds act upon one another, and what products are formed. The question of composition is one of the first that presents itself, and this must be studied before other questions can be intelligently discussed. What, then, are the most prominent facts which we have learned in studying the elements and com- pounds which have thus far been taken up ? Kinds of Chemical Reactions. — All chemical reactions may be classified under three heads : (1) Those which consist in direct combination; (2) Those which consist in direct decomposition; and (3) Those which involve the interaction of two or more elements or compounds and the formation of two or more compounds. This is known as double decomposition or metathesis. Direct Combination. — We have had to deal with a num- ber,of examples of each of these kinds of reactions. As 4°<> COLLEGE CHEMISTRY. examples of the first kind already studied the following may be mentioned : The combination of hydrogen and chlorine to form hydrochloric acid; the formation of ammonium chloride from ammonia and hydrochloric acid; the formation of calcium hydroxide from calcium oxide and water; the formation of nitrogen peroxide from nitric oxide and oxygen; and the formation of carbon disulphide from carbon and sulphur. As regards the combination of hydrogen and chlorine, it should be remarked that this act is the same in principle as that of metathesis. Strictly speaking, it is not a case of direct combination, as we understand it. For, as we have seen, according to the molecular theory, free chlorine and free hydrogen consist of molecules which are made up of two atoms each. Therefore, when these elements are brought together the molecules are first decomposed into atoms before the act of union can take place. The two acts are represented by the two equations following : Cl 2 + H a = CI + CI + H + H; CI + CI + H + H = 2HC1. In the Case also of the union of hydrochloric acid and ammonia it appears probable that a serious disarrangement of the constituent atoms is necessary in order that the act of combination may take place. According to the am- monium theory, ammonium chloride is represented by the f H H formula N I H, which means that the atom of chlorine H CI and four atoms of hydrogen are in combination with the atom of nitrogen. But in order that a compound of this constitution may be formed from ammonia and hydro- chloric acid, it is necessary that the molecule of hydro- chloric acid should be broken down into its constituent atoms. So that this case of apparent direct combination DIRECT DECOMPOSITION. 4°7 is, as far as we can judge, in reality more complicated than it appears, and should be represented by the two equations : NH 3 + HC1 = NH 3 + H + CI; NH 3 + H + CI = NH 4 C1. All other cases of apparent direct combination are prob- ably of the same character, so that it is doubtful whether a single case of simple direct combination is known. Direct Decomposition. — As examples of direct decom- position the following cases may be cited : The decomposition of mercuric oxide by means of heat into mercury and oxygen; that of ammonium chloride into ammonia and hydrochloric acid by heat; that of potassium nitrate into potassium nitrite and oxygen by heat; that of phosphorus pentachloride into the trichloride and chlorine by heat; that of ammonia into hydrogen and nitrogen by continued action of electric sparks; and that of nitrogen iodide by contact with a solid substance. On close examination of each of the above cases, which are fairly typical and as simple as any that could be chosen, it will be seen that no one of them is merely a case of decomposition; for even though we must assume that the first result in each case is the setting free of the atoms of one or two elements, we must also assume that these atoms unite again to form other molecules either of ele- ments or compounds. Thus, when mercuric oxide is decomposed we get mercury and oxygen. As far as can be determined, the mercury atoms do not unite with each other, but the oxygen atoms do, so that the total action involves decomposition and afterwards combination as represented in the equations 2HgO = Hg + Hg + + 0; Hg + Hg + + = Hg + Hg + O a . In the case of the pentachloride of phosphorus, it is prob- able that the two atoms of chlorine are first given off from each molecule of the chloride, leaving a molecule of the 4°» COLLEGE CHEMISTRY. trichloride, but the atoms of chlorine afterwards unite to form molecules as represented thus : PC1 6 = PC1 8 + CI + 01; PC1 3 + 01 + CI = PC1 S + CI,. Similar statements hold good for all other cases of direct decomposition. Metathesis. — This is the most common kind of chem- ical action, and indeed from what has been said in regard to direct combination and direct decomposition it will be seen that there is no essential difference between them and metathesis. Most of the reactions with which we have had to deal are examples of double decomposition or metathesis, as: The formation of salts by the action of bases upon acids; the formation of the sulphides of arsenic, antimony, and bismuth by the action of hydrogen sulphide upon solutions of compounds of these elements; the setting free of hydrochloric and nitric acids by the action of sulphuric acid upon chlorides and nitrates; of carbon dioxide and oxides of nitrogen by the action of acids upon carbonates and nitrites; and of ammonia by treating am- monium salts with lime. As simple an example of this kind of action as can be given is that of the formation of hydrogen and potassium chloride from potassium and hydrochloric acid gas. The molecular weight of potassium is not positively known, but, assuming its molecule to be made up of two atoms, the action must be represented in this way : K, + 2HC1 = 2K01 + H„. The next stage of complication is exhibited in the reaction following: KI + HC1 = KC1 + HI. Examples similar to the latter, but somewhat more com- plicated, are these: 2KOH + H 2 S0 4 = K 2 S0 4 + H 2 0; CaCl 2 + H 2 S0 4 = CaS0 4 + 2H0L THE CAUSE OF CHEMICAL REACTIONS. 4°9 The Cause of Chemical Reactions The prime cause of chemical reactions is something which we think of as an attractive force exerted in different degrees between the different elements. When any elements or compounds are brought together under certain conditions the tendency is always towards the formation of the most stable com- pounds of those elements which can be formed under the given conditions. Thus, potassium sulphate and water are more stable forms of combination of the elements hydrogen and oxygen, and potassium, sulphur and oxygen, than sulphuric acid and potassium hydroxide are under the conditions under which the action takes place. So also the system composed of potassium chloride and hydriodic acid is more stable than that composed of potassium iodide and hydrochloric acid under the conditions of the action. "Why the one system is more stable than the other we do not know, for we do not know what relations exist between the atoms in the molecules. It is convenient to think of that which causes the atoms to unite to form compounds as an attractive force. It is evident that this force is more strongly exerted between some elements than between others. That between chlorine and hydrogen is, for ex- ample, much stronger than that between chlorine and nitrogen or oxygen. Owing, however, to the complicated character of most chemical reactions, it is extremely diffi- cult to make measurements of this force. An Ideal Chemical Reaction — In every case in which two compounds act upon each other to form two new ones, several forces must be at work, as we have seen. Suppose, for example, AB and CD act upon each other in the gaseous condition to form two compounds BC and AD, also both gaseous. The normal course of such a reaction would lead to the formation not only of the two compounds Z?Cand AB, but AB and CD would also be present in the resulting system. For A has the power to combine with B as well as with D, and C has the power to combine with D as well as with B. In the system we should have operating the tendency of A to combine with B, and that 41° COLLEGE CHEMISTRY. of A to combine with D; that of to combine with D, and that of C to combine with B. As these forces operate simultaneously, equilibrium is established when certain quantities of the four possible compounds are formed, the quantities depending in the first instance upon the relative strengths of the forces at work. It is difficult to give an illustration of this kind of rela- tion, because the study of chemical changes between gases that give rise to the formation of gaseous products is diffi- cult, and the study of changes in solution is complicated by the action of the solvent and, if the solvent is water and the substances are acids, bases, or salts, this action plays a controlling part. Influence of Mass. — When hydrogen sulphide is passed into a solution of cadmium chloride, cadmium sulphide is precipitated, the reaction being represented by the equa- tion (1) CdCl, + H 2 S = CdS + 2HC1. On the other hand, when hydrochloric acid is added to cadmium sulphide the latter dissolves and the action is represented by the equation (2) CdS + 2HC1 = CdC] 2 + H 2 S. This action is plainly the reverse of that which takes place when hydrogen sulphide acts upon cadmium chloride in solution. The difference between the two cases is due to the relative masses of hydrogen sulphide and of hydro- chloric acid that are brought into action. If the hydro- chloric acid is very dilute and the hydrogen sulphide is present in excess, reaction (1) takes place and is complete. If the conditions are reversed, then reaction (2) takes place. If considerable hydrochloric acid is added to a solution of cadmium chloride, and hydrogen sulphide passed into the solution, all of the cadmium will not bo precipitated as sulphide, and the amount precipitated will depend upon the amount of hydrochloric acid present. INFLUENCE OF MASS. 4" A reaction that can take place in two directions or that is reversible is commonly represented thus: CdCl a + H,S £ CdS + 2HC1. Equilibrium is established in such cases under definite conditions, and the law governing the equilibrium has been discovered. Guldberg and Waage were the first ones to formulate the law of mass action. The law is as follows : When two substances act upon each other, the action is proportional to the active mass of the substances taking part in the change. By active mass is meant the molecular concentration, or, in other words, the number of gram-molecules per .litre, a gram-molecule being the number of grams corre- sponding to the molecular weight of the substance. A gram -molecule of sodium chloride, for example, is 58.5 grams, etc. The following quotation from the memoir by Guldberg and Waage will make the subject somewhat clearer: "When two substances A and B are transformed into two new substances A' and B', the chemical force with which A and B act upon each other is measured by the quantity of the new substances formed in unit time. " The quantity of a substance in unit volume of the compound in which the chemical change takes place we call the 'active mass of the substance. " The chemical force with which two substances A and B act upon each other is equal to the product of their active masses, multiplied by the coefficients of affinity. "By coefficient of affinity a coefficient is understood which is dependent upon the chemical nature of the two substances and upon the temperature. If the active masses of A and B are represented by p and q, and the coefficient of affinity by h, then the force acting between A and B is expressed by Tcpq. " When in a chemical process A and B are transformed 412 COLLEGE CHEMISTRY. into A' and B', and A' and B' can at the same time be transformed into A and B, equilibrium will be established when the force acting between A and B is equal to the force acting between A' and B'. " If the active masses of A' and B' are represented by p' and q' and their coefficient of affinity by k', the chemical force: acting between A' and B' is expressed by k'p'q'. "The condition of equilibrium is therefore expressed by the equation kpq = k'p'q'." Reactions may be Complete if one of the Products Formed is Insoluble or Volatile. — When two substances which by interaction can form an insoluble product are brought together, the reaction generally takes place and is complete. When the substances are brought together we may imagine that, owing to interaction, a small quan> tity of the insoluble compound is formed at once. If this product were soluble, the action would stop before it is complete, because this new product would itself exert its action upon the system. Being insoluble, however, it is removed from the sphere of action, and the same reaction which caused the formation of the first particles of it can now be repeated, and so on, until the reaction is complete. This is illustrated in the action of sulphuric acid upon barium chloride in solution. The two substances react as represented in this equation : BaCI + H a S0 4 + Aq = BaS0 4 +• HC1 + Aq. The symbol Aq is simply intended to indicate that the reaction takes place in solution. If barium sulphate were soluble, all four substances — barium chloride, sulphuric acid, barium sulphate, and hydrochloric acid — would be present in the solution after the establishment of equili- brium. But, being insoluble, it is removed, and new quantities are formed as long as the substances necessary for its formation are present in the solution; that is, until either all the barium chloride is decomposed or all the sul- phuric acid is removed. Eeactions involving the formation Complete reactions. 413 of insoluble compounds or precipitates are among the most common with which we have to deal, particularly in the various operations of analytical chemistry. Again, when two substances which can form a volatile product are brought together the reaction generally takes place and is complete. The reason why a reaction of this kind is complete is the same as that given in the case of the formation of an insoluble compound. Each successive portion of the volatile product formed is removed, and the reaction which gave rise to it proceeds as long as the necessary substances are present. This kind of action has been repeatedly illustrated. It is that, for example, which is seen in the liberation of hydrochloric acid from a chlo- ride by the action of sulphuric acid; of carbon dioxide by the action of an acid upon a carbonate; and of ammonia by the action of lime upon ammonium chloride. An interesting example of the combined influence of mass and the volatility of the product is seen in the action of heated iron upon an excess of steam, and of the oxide of iron upon an excess of hydrogen. When steam is passed over heated iron, action takes place thus : 4H 2 + 3Fe = Fe 3 4 + 4H 2 . Hydrogen is liberated and the oxide of iron formed. When, however, hydrogen is passed over heated oxide of iron the reverse reaction takes place: Fe 3 4 + 4H 2 = 3Fe + 4H 2 0. Owing to the excess of steam always present in the first reaction, hydrogen is constantly formed and constantly being removed. Undoubtedly the hydrogen formed acts to some extent upon the oxide, but the other reaction always takes place to a greater extent. The opposite is true when the oxide is heated in an excess of hydrogen. The principal reaction which takes place in this case is that of the hydrogen upon the oxide of iron, and the steam is carried out of the field almost as soon as formed, so that the reduction of the oxide of iron continues. 414 COLLEGE CHEMISTRY. A, Dissociation. — It has already been stated (p. 87) that T when water-vapor is heated to a sufficiently high tempera- ture it is decomposed into the two gases hydrogen and oxygen. This decomposition does not take place suddenly when a certain temperature is reached, but it begins at a comparatively low temperature and increases as the tem- perature is raised. For any given temperature and pressure the amount of decomposition is definite. On cooling a mixture of water-vapor and hydrogen and oxygen, the two elements combine to a greater and greater extent as the temperature is lowered. A condition of equilibrium is established under each new set of condi- tions. The gradual decomposition of a substance into its constituents which reunite when the temperature is lowered is called dissociation. All cases of dissociation are examples of reversible reactions, and they also illustrate the phenomena of equilibrium. Other examples besides that mentioned are the change of nitrogen peroxide from the form N 2 4 to that of NO a . This may be represented thus: NA ~ NO, + N0 2 . The dissociation of hydriodic acid into hydrogen and iodine may also be mentioned: In all of these cases the substances taking part are gaseous. When, however, calcium carbonate breaks down into lime and carbon dioxide two of the three substances are solids and one is a gas : CaC0 3 ^CaO + C0 2 . Again, the dissociation of ammonium chloride into am- monia and hydrochloric acid involves the interaction of three substances, two of which are gases and one a solid : NH 4 C1 £ NH a + HCL DISSOCIATION. 41 S It would lead too far to discuss these reactions in detail here. The explanation of the phenomena of dissociation of gases is found in the kinetic theory of gases. According to this theory, the molecules of a gas at a given tempera- ture are moving with different velocities, though the average velocity of all the molecules is the same at the same temperature. Now, it is highly probable that the motion of the atoms within the molecules partakes of that of the molecules themselves, so that the motion of the atoms in the molecules with the greatest velocity is prob- ably the greatest, and, in these, decomposition will take place first. When a compound gas is heated, we can easily conceive that even at a comparatively low temperature the motion of some of the molecules will be sufficient to cause their decomposition, and, as the average motion of all the molecules is constant for a given temperature, the amount of decomposition will be constant for that temperature. A& the molecules are, however, moving in every direction and constantly colliding, a molecule which is decomposed at one instant may be re-formed at the next, and one that is not deeomposed may acquire motion enough to cause its decomposition. Though, as is believed, these changes are constantly taking place at every temperature, still, as has been said, the number of molecules which will be decom- posed in a given mass at a given temperature and pressure will always be the same. The higher the temperature, then, the greater the number of molecules in the condi- tions which cause decomposition, and the smaller the number of those in the conditions favorable to formation. At each temperature and pressure an equilibrium is estab- lished, the . number of molecules decomposed being equal to the number formed. It is obvious that, if one of the products of decomposition is removed, the conditions are entirely changed. Then the possibility of recombination will not exist, and total decomposition can be effected at a lower temperature than that required for total decom- position in the process of dissociation proper. 4 1 6 COLLEGE CHEMtSTRY. Electrolysis.— Some chemical compounds in solution in water conduct electricity, and at the same time they undergo decomposition. Thug, hydrochloric acid in solu- tion in water conducts electricity, and its constituents hydrogen and chlorine are obtained, the hydrogen appear- ing at the negative and the chlorine at the positive pole. Compounds that act in this way are called electrolytes. When a current of electricity acts upon solutions of differ- ent salts, equivalent quantities of the metals are deposited by the same current in the same time. This is Faraday's Law. Thus if the same current were passed simul- taneously through solutions of silver nitrate, AgNO s , mercuric nitrate, Hg(N0 3 ) 2 , cupric sulphate, CuS0 4 , and ferric chloride, FeCl 3 , it would be found that for every 107.93 parts by weight of silver deposited there would be 100.15 parts by weight of mercury deposited, 31.8 of copper, and 18.67 of iron. These are equivalent quantities of these metals — quantities that take the place of one part by weight of hydrogen — and are to be distinguished from atomic quantities. Those elements which appear at the negative pole, or cathode, are called electro-positive, and those which appear at the positive pole, or anode, are called electro-negative. Those elements which we call acid-forming are electro-negative, while hydrogen and the base-forming elements are electro-positive. The ions that tend towards the positive pole, or anode (the acid-form- ing, or electro-negative elements, etc.), are the anions, while those which tend towards the negative pole, or cathode (base-forming, or electro-positive elements, etc.), are the cations. The electrolysis of chemical compounds is not generally a simple decomposition into two constituents. Thus, when copper sulphate, CuS0 4 , is decomposed, the copper is deposited at the negative pole; but no such com- pound as S0 4 appears at the positive pole. This, if formed, forms sulphuric acid and gives off oxygen. Both oxygen and sulphuric acid as a matter of fact appear at the positive pole. The changes involved may be represented thus : CuS0 4 = Ou + S0 4 ; S0 4 + H 2 = H 2 S0 4 + 0. ELECTROLYTIC DISSOCIATION. 417 / Electrolytic Dissociation.— -It has been known for a long time that a very weak electric current acting upon a solu- tion of an electrolyte is sufficient to cause the ions to appear at the poles. This fact is inexplicable if it is assumed that the current is the cause of the decomposition of the electrolyte. This, and some other facts which will be referred to farther on, make it probable that electrolytes are at least to some extent decomposed into their constit- uent ions when they are dissolved in water; that these ions charged with electricity transfer their charges in the solu- tion and thus conduct the current; and that when an ion charged with negative electricity reaches the positive pole its electricity is discharged, and the ion then ceases to be an ion and becomes an element in the free state or some compound which appears either as such or in the form of other products. According to this conception, the act of solution of an electrolyte, in water at least, involves partial breaking down or dissociation of the compound into its ions. The extent of this breaking down is determined primarily by the concentration of the solution — the greater the dilution the greater the dissociation. At infinite dilu- tion there is complete dissociation. A water solution of hydrochloric acid containing 36.18 grams of the acid in 1000 litres has been shown to be completely dissociated, or it is to be regarded as containing hydrogen ions and chlorine ions. These and all other ions are carefully to be distinguished from the atoms or definite compounds. An ion always carries with it a certain charge of elec- tricity. When this is discharged the ion becomes either an element or a compound in the free state. When a solution of one electrolyte acts upon a solution of another the reaction observed is probably due to the interaction of the ions, and it is farther probable that, so far as the compounds are present in the undissociated condition, they do not act upon each other. If this vie-w be correct the reactions most familiar to us are reactions of ions, and not of elements or compounds. When, for example, an acid acts upon a base in solution it appears that, so far as 4i8 COLLEGE CHEMISTRY. they react, they are in dissociated condition. Thus hydro- chloric acid and sodium hydroxide are to be regarded as acting as represented in the following equation : H + CI + Na + OH = Na + CI + H 2 0. The act consists in the union of the hydroxyl ion of the base with the hydrogen ion of the acid to form water, the sodium and chlorine ions remaining as ions as in a dilute solution of sodium chloride. In the case of nitric acid and potassium hydroxide the following equation represents the reaction at infinite dilution: H + NO, + K + OH = K + N0 3 + H,0. And so also whenever the act of neutralization takes place there is simply a union of hydrogen ions with hydroxyl ions to form water. This conception finds strong confirmation in the fact that the heat evolved in neutralizing equivalent quantities of all acids at infinite dilution is always the same — a fact difficult to explain if it is assumed that in the act of neutralization a salt is formed in the solution. In the following table the heats of neutralization of a few acids and bases are given : Acid and Base. Heat of Neutral. Hydrochloric acid and sodium hydroxide. . . . 13,700 " " " lithium " . . 13,700 " " " potassium " . . 13,700 " " " barium " . . 13,800 " " " calcium " .. 13,900 Hydrobromic " " sodium " . . 13,700 Nitric " " " " .. 13,700 Iodic " " " " . . 13,700 As will be seen, the heat of neutralization is the same no matter what the base or what the acid may be, and as has been pointed out this fact is easily understood, if the ELECTROLYTIC DISSOCIATION. 4*9 act of neutralization consists in the union of a hydroxyl ion with a hydrogen ion to form water. If the strength of an acid is determined by the extent to which it is dissociated, then of course those acids that are most readily dissociated are the strongest. By every method available hydrochloric and nitric acids are found to be the most readily dissociated and they are the strongest acids. The electrical method for the determina- tion of the strength of acids is based upon this theory. The applicability of this theory to the explanation of the most common reactions that take place in solution is at present attracting the attention of chemists. Those reac- tions which are made use of for the purpose of detecting the presence of the various elements appear in fact, as has already been stated, to be reactions of the ions, and when these ions are not present the reactions are not observed. For example, when silver nitrate in solution is added to sodium chloride in solution a precipitate of silver chloride is formed, the reaction taking place as represented in this equation : AgNO, + NaCl = AgCl + NaN0 3 . This reaction seems, however, to be due to the fact that ions of silver and of chlorine are present. These coming together form the insoluble molecules silver chloride which is then precipitated : Ag + NO, + &a + CI = AgCl + Na + N0 S . There are many compounds that contain chlorine and yet do not give a precipitate of silver chloride when treated with silver nitrate. It is believed that in these cases the compound is not ionised by the solvent in such a way as Qp to yield ions of chlorine, and that therefore there areVchlorine ions present. An example of a compound that contains chlorine and yet does not give the reac- tions of this element is potassium chlorate, KC10 3 . A solution of this salt does not give a precipitate with silver 42° COLLEGE CHEMISTRY. nitrate. It is probable that the reason of this is that the salt gives the ions K and C10 s , the latter acting quite differently from the ion 01, as we should naturally expect. Definition of Acids and Bases in Terms of the Theory of Electrolytic Dissociation.— According to what has just been said it appears that when any acid is dissolved in water it is dissociated into hydrogen ions and something else. This power to yield hydrogen ions is characteristic of acids. On the other hand, the most marked examples of bases, such as sodium and potassium hydroxides, when dissolved in water give hydroxyl ions, and this conduct is characteristic of bases. * Raoult's Methods for the Determination of Molecular Weights. — One great difficulty encountered in the study of chemical compounds is the determination of the molec- /ular weights of those which are not gases or cannot be converted into vapor by heat. From some studies on the freezing-points of solutions, it appears that quantities of compounds proportional to their molecular weights cause the same lowering of the freezing-points, provided the solvent does not act chemically upon the compound. This fact makes it possible to determine the molecular weights of substances which cannot be converted into vapor, but which can be dissolved. The application of the me 1 '•.od is simple. Suppose water to be the solvent used. "We know that this liquid solidifies or freezes at 0°. Now, it is found that by dissolving a certain quantity of some sub- stance in a certain quantity of water the freezing-point is lowered say .5°. Further, the quantities of other sub- stances which are necessary to lower the freezing-point of the same quantity of water to the same extent can be determined. These quantities are proportional to the molecular weights according to the law of Eaoult. If, therefore, among the substances studied there is one th.3 molecular weight of which can be determined by the method of Avogadro, it is possible to determine A he molec- ular weights of all of them by the method of Eaoult, as will readily be seen. * , \ DISSOCIATION OF A DISSOLVED^ SUBSTANCE. 4™ So, also, it has been shown that quantities of compounds proportional to their molecular weights cause the same raising of the boiling-points, proYided the solvent does not act chemically upon the compounds or cause them to break down into their ions. Convenient methods have been devised for the determination of molecular weights of dip solved substances, the methods being based upon observa- tions on the boiling-points and freezing-points. It should be noted that, when the molecular weight of a substance in solution has been determined, it does not follow that the substance has the same molecular weight when in the solid condition. This is a matter in regard to which we have little knowledge. It is quite possible that the molecules of solid substances may be made up of large aggregates of the simple molecules that probably exist in solutions or in vapors. There is, however, no method at present known that makes a determination of the com- plexity of these molecules, or molecular aggregates,, possible. Determination of the Extent of Dissociation of a Dis- solved Substance. — The effect upon the boiling-point or freezing-point of a solution caused by the presence of a dissolved substance is proportional to the number of mole- cules in the solution or the number of individual particles^ whether these are undecomposed molecules or the ions formed as a result of dissociation. Any substance that is dissociated in solution will give abnormal results if the attempt is made to determine its molecular weight by observations on the boiling-point or the freezing-point of its solutions. This method is therefore not applicable to solutions of electrolytes. On the other hand, the study of such solutions has shown that there is an increased lowering of the freezing-point for the same weight of solvent as the dilution becomes greater, a fact that points ■ clearly to the conclusion that as the solution is diluted' there is greater and greater dissociation, and advantage can be taken of this fact for the purpose of determining 422 COLLEGE CHEMISTRY. the extent to which dissociation has taken place in a solu- tion of an electrolyte in water. In general only organic compounds come within range of the methods of Raoult. These methods are now exten- sively used in the study of the compounds of carbon, simple forms of apparatus having been devised for this purpose. The recognition of the fact that electrolytes do not obey the simple law that holds good in the case of non-electrolytes led Arrhenius to the idea that the former are dissociated in solution — an idea which has proved of great service to the science, and is likely to revolutionize the views of chemists in regard to the action of chemical substances upon each other in solution. Osmotic Pressure. — During the last century it was observed that, when a vessel is filled with alcohol and the vessel "tightly covered by tying a bladder over its mouth, and the whole then immersed in water, the bladder is stretched outward, showing that the liquid from without has found its way into the vessel through the bladder. For many years investigations on phenomena of this kind were carried on, but no results of a general character of special importance from the chemical point of view were reached until recently. In 1877 Pfeffer, the botanist, gave an account of a large number of experiments, and laid the foundation of our present knowledge of the laws of osmotic phenomena. The simplest phenomenon of this kind is seen when a wide glass tube, tightly closed with a piece of bladder at the lower end, is partly filled with alcohol and then placed in a vessel of water, so that the level of the liquids inside and outside the tube is the same. If the tube is fixed in this position, the liquid will soon be found to rise in the tube. There is pressure from without inward. This is called osmotic pressure. Pfeffer's inves- tigations had to do with the measurement of the pressure exerted by different substances under different conditions of temperature and dilution. Instead of using a bladder, which is not capable of much resistance, he used mem- branes made by precipitating copper ferrocyanide in the SPECIFIC HEAT AND ATOMIC WEIGHTS. 42 3 pores of clay cells. Somewhat later Van't Hoff showed that the results obtained by Pfeffer lead to the following remarkable laws governing osmotic pressure : 1. The pressure is proportional to the concentration, or it is inversely proportional to the volume in which a definite quantity of the dissolved substance is contained. 2. The pressure increases for constant volume, propor- tionally to the absolute temperature. 3. Quantities of dissolved substances which are in the ratio of the molecular weights of these substances exert equal/ pressures at the same temperatures. / These laws are analogous to the well-known laws of gas- pressure, the third being plainly analogous to the law of Avogadro. Another form of the third law, together with an extension of it, is this : Dissolved substances exert the same pressure, in the form of osmotic pressure, as they would exert were they gasified, at the same temperature, without change of volume. Notwithstanding the simplicity of this law, no practical method for determining molecular weights based upon it has yet been devised. The difficulties are, however, of an experimental nature. Relations between Specific Heat and Atomic Weights. — The fact that there is a method for the determination of atomic weights founded upon the relations existing between these weights, and the specific heat of the ele- ments, has been mentioned. It has been found that, when equal weights of different elements are exposed to exactly the same source of heat, they require different lengths of time to become heated to the same temperature. Given the same heating power, it requires 32 times as long to raise the temperature of a pound of water 10, 20, or 30 degrees as it does to raise the temperature of a pound of mercury the same number of degrees; or it takes 32 times as much heat to raise a pound of water 10, 20, or 30 degrees as it does to raise a pound of mercury the same number of degrees. Starting at the same temperature, the quantity of heat required to raise the temperature of a 424 COLLEGE CHEMISTRY. certain weight of a substance one degree, . as compared with the quantity of heat required to raise the temperature of the same weight of water one degree, is called the specific heat of the substance. Thus, from what was said above, the specific heat of mercury is g 1 ^, or, in decimals, 0.03125. In a similar way it can be shown that the specific heat of gold is 0.03244; of zinc, 0.0955; of silver, 0.057; of copper, 0.0952. Now, when solid elements are examined with reference to their specific heats, a very simple relation is found to exist between the numbers expressing the specific heats and the atomic weights. This relation will be made clear by a consideration of a few cases : Element. Specific Heat. Atomic Weight. Silver 0.0570 107.93 Zinc 0.0955 65.4 Cadmium 0.0567 112.4 Copper 0.0952 63.6 Tin 0.0562 118.5 An examination of this table will show that the atomic weights are inversely proportional to the specific heats. We have 107.93 112.4 107.93 65.4 63.6 118.5 0.0955 : 0.0570; 0.0952:0.0567; 0.0562 : 0.0570; etc. These proportions are only approximately correct; but it must be remembered that the means for the determina- tion of atomic weights and specific heats are not perfect, and in both sets of figures there are undoubtedly small errors. Hence such slight variations from absolute agree- ment in these proportions should occasion no surprise. The agreement is sufficient to indicate a close connec- tion between the two sets of figures. This connection may be stated in another way : The product of the atomic SPECIFIC HEAT AND ATOMIC WEIGHTS. 425 weight by the specific heat is a constant. Thus, in the above cases: 107.93 X 0.057 = 6.16 65.4 X 0.0955 = 6.24 112.4 X 0.0567 = 6.37 63.6 X 0.0952 = 6.05 118.5 X 0.0562 = 6.66. From the above it appears that the quantity of neat necessary to raise masses of the elements proportional to their atomic weights the same number of degrees is the same in all cases. Suppose two elements to have the atomic weights 2 and 4. Their specific heats would be to each other as 2 to 1. That is to say, it would require twice as much heat to raise the temperature of a given mass of the element with the atomic weight 2 a certain number of degrees, as it would require to raise the tem- perature of the same mass of the element with the atomic weight 4 the same number of degrees. But to raise the temperature of masses of these two elements proportional to their atomic weights would require the same quantity of heat. This fact may be stated thus : The atoms of all elements have the same capacity for heat. This is only another way of stating that, to raise the tem- perature of an atom one degree, the same quantity of heat is always necessary. Now, if we assume that the constant obtained by mul- tiplying the specific heats by the atomic weights is 6.4, which is about the average of the different values found, then it is plain that, if we divide this number by the specific heat of an element, we shall obtain a number which is very near the atomic weight. If we call the atomic weight A, and the specific heat H, the following equation expresses the relation : A- H . If this law is without exceptions, it is plain that, in order 426 COLLEGE CHEMISTRY. to determine the atomic weight of an element, it is only necessary to determine its specific heat, and divide this into 6.4. The result will be very nearly the atomic weight. Knowing thus very nearly what the atomic weight is, it is a comparatively simple matter to determine it with great accuracy by means of chemical analysis. Unfor- tunately there are some marked exceptions to the law. Exceptions to the Law of Specific Heats.— The elements glucinum, carbon, boron, and silicon form exceptions to the law of specific heats as this law has been stated above. At ordinary temperatures they do not follow the law. As the temperature is raised, however, the specific heat of these elements changes markedly, until finally, in the cases of carbon and silicon, a point is reached beyond which there is no marked change. Thus, at 600° the specific heat of diamond is 0.441, and at 985° it is 0.449. That of silicon is 0.201 at 185°, and 0.203 at 332°. At these temperatures the elements obey the law. Prom elaborate studies which have been made on this subject, it appears that the law should be modified to read as follows: The specific heats of the elements vary with the tem- perature; but for every element there is a temperature, T, above which variations are very slight. The product of the atomic weight by the constant value of the specific heat is nearly a constant, lying between 5.5 and 6.5. Notwithstanding the irregularities referred to, the law of specific heats, commonly called, from the discoverers, the law of Dulong and Petit, is of great value in the determination of atomic weights. N, CHAPTER XXIV. BASE-FORMING ELEMENTS.— GENERAL CONSIDERA- TIONS. Introductory. — The elements thus far studied belong for the most paft to the class of acid-forming elements, or those whose compounds with oxygen and hydrogen have acid properties. All the members of Family VII, Group B, are acid-forming, while the single member of Group A of the same family is both acid-forming and base-forming. All the members of Family VI, Group B, are acid-forming, while the members of Group A of this family are both acid-forming and base-forming. In Family V, Group B, a gradation of properties is observed, the group beginning with strongly marked acid-forming elements and end- ing with an element, bismuth, which is more basic than acid in character. The elements of Group A, Family V, are both acid-forming and base-forming, but they have not as sharply marked characteristics as the elements of Families VI and VII. Passing now to Family IV, it was found that the two most important members, carbon and silicon, belong to Group A. These two elements always act as acid-formers. A gradation of properties is observed in passing from silicon to thorium. The members of Group B of this family have the properties of the base- forming elements much more strongly marked than those of the acid-formers. There are still four families to be studied. These are Families I, II, III, and VIII, the members of which are almost exclusively base-forming elements. The compounds of these elements with hy- drogen and oxygen are bases, or, in other words, have the 427 428 COLLEGE CHEMISTRY. power to neutralize acids. Their oxides are for the most part basic. An exception to this is found in the case of boron, already treated of, which forms a weak acid — boric acid. Its oxide is only slightly basic. The most. strongly marked examples of base-forming elements are those which occur in Family J, Group A; then follow in order those of Group A, Family II, and Group A, Family III. The resemblance between the members of Group B, Family I, and those of Group A of the same family is less striking than the resemblance between the two groups of any other family. Between the members of Group B, Family II, and those of Group A of the same family there is a general resemblance, while there are also differences. A similar remark applies to the relations between Groups A and B, Family III. The members of Family VIII occupy a some- what exceptional position, as has already been pointed out. Each group of which this family consists is made up of three very similar elements with atomic weights which differ but little from one another. Metallic Properties. — It has long been customary to divide the elements into two classes — the metals and the non-metals. This classification was originally based upon differences in the physical properties of the elements, the name metal being applied to those elements which have what is known as a metallic lustre, are opaque, and are good conductors of heat and electricity. All those ele- ments which do not have these properties are called non- metals. Gradually the name metal came to signify an element which has the power to take the place of the hydrogen of acids and form salts, and the name non-metal to signify an element which has not this power. This classification, as will be seen, is practically the same as that which divides the elements into acid-forming and base-forming. The latter are the metals, the former are the non-metals. The imperfection of this classification has already been commented upon, the imperfection arising from the fact that some elements belong to both classes. STUDY OF THE BASE-FORMING ELEMENTS. 4*9 Order in which the Base-forming Elements will be Taken Up. — In studying the base-forming elements, it appears best to begin with those which have the most strongly marked character. These are the members of Family I, Group A. It further appears best to adhere as closely as possible to the arrangement of the periodic system. Accordingly, the following order will be observed in the presentation of the elements yet to be studied : 1. Elements of Family I, Group A, or the Potassium Group, consisting of lithium, sodium, potassium, ru- bidium, and caesium. 2. Elements of Family II, Group A, or the Cglciwai Group, consisting of glucinum, magnesium, calcium, strontium, barium, and erbium. 3. Elements of Family III, Group A, or the Aluminium Group, . consisting of aluminium, scandium, yttrium, lan- thanum, and ytterbium. 4. Elements of Family I, Group B, or the Copper Group, consisting of copper, silver, and gold. 5. Elements of Family II, Group B, or the Zinc Group, consisting of zinc, cadmium, and mercury. , „ , 6. Elements of Family III, Group B, or the Gallium Group, consisting of gallium, indium, and thallium. 7. Elements of Family IV, Group B, or the Tin Group, consisting of germanium, tin, and lead. 8. Elements of Family V, Group A, or the Vanadium Group, consisting of vanadium, columbium, didymium, and tantalum. 9. Elements of Family VI, Group A, or the Chromium Group, consisting of chromium, molybdenum, tungsten, and uranium. 10. Elements of Family VII, Group A, or the Man- ganese Group, of which manganese is the only representa- tive. 11. Elements of Family VIII, of which there are three , , groups: \ (A) The Iron Group, consisting of iron, nickel, and cobalt j 43° COLLEGE CHEMISTRY. (B) The Palladium Group, consisting of ruthenium, rhodium, and palladium; and (0) The Plalinum Group, consisting of osmium, irid- ium, and platinum. Occurrence of the Metals — One of the first questions that suggests itself in connection with each element is, In what forms of combination does it occur in nature ? The chemical compounds which occur ready-formed in nature are called minerals; and the minerals, and mixtures of minerals, from which the metals are extracted for practical purposes are called ores. The most common ores are oxides and sulphides. Examples of these are the ores of iron, tin, copper, lead, and zinc. The carbonates also occur in large quantity in nature, and are used for the purpose of preparing some of the metals. The carbonate of zinc, for example, is a valuable ore of zinc. Extraction of the Metals from their Ores. — The detailed study of the methods used in the extraction of the metals from their ores is the object of metallurgy. Besides the methods used on the large scale, there are others which are only used in the laboratory, The most common method of extracting metals from their ores is that used in the case of iron, which consists in heating the oxides with charcoal. If the ores used are not oxides, they must first be converted into oxides before this method is appli- cable. This can generally be accomplished by heating the ores in contact with the air. Under these circumstances the natural carbonates, sulphides, and hydroxides, are converted into oxides. These changes are illustrated by the following equations : FeC0 3 = FeO + C0 2 ; 2Fe0 + = Fe 2 3 ; 2FeS 2 + 110 = Fe 2 3 + 4SO a ; 2Fe(0H) 3 = Fe 2 3 + 3H 2 0. By heating metallic oxides with carbon at the high tem- peratures attainable by means of powerful electric currents THE PROPERTIES OF THE METALS. 431 in the electric furnace some oxides are easily reduced that cannot be reduced by other means. A second method consists in reducing the oxide by heat- ing it in a current of hydrogen. This has been illustrated in the action of hydrogen upon copper oxide, when the following reaction takes place : CuO + H 2 = H 2 + Cu. The method is efficient for many oxides, but is expensive and is not used on the large scale. Another method of extraction consists in treating the chloride of a metal with sodium. This is illustrated in the preparation of magnesium, which is made by heating together magnesium chloride and sodium : MgCL. + 2Na = 2NaCl + Mg. Such a method is employed only in case it is impossible or extremely difficult to reduce the oxide. Besides the above methods, there are others which will be described under the individual metals. The Properties of the Metals.— As we shall find, the metals differ very markedly from one another. Some are light, floating on water, as lithium, sodium, etc. ; some are extremely heavy, as lead, platinum, etc. Some combine with oxygen with great energy; others form very unstable compounds with oxygen. Some form strong bases; others form weak bases. In general, those elements which are lightest, or which have the lowest specific gravity, are the most active chemically, while those which have the highest specific gravity are the least active. Among the former are lithium, sodium, and potassium; among the latter are lead, gold, and platinum. Compounds of the Metals — The principal compounds of the metals may be conveniently classified as : a. Compounds with fluorine, chlorine, bromine, and iodine; or the fluorides, chlorides, Iromides, and iodides. 43 2 COLLEGE CHEMISTRY. b. Compounds with oxygen, and with oxygen and hydrogen; or the oxides and hydroxides. c. Compounds with sulphur, and with sulphur and hydrogen; or the sulphides and hydrosulphides. d. Compounds with nitrogen; or the nitrides. e. Compounds with carbon and with silicon; or the carbides and silicides. f. Compounds with the acids of nitrogen; or the nitrates and nitrites. g. Compounds with the acids of chlorine, bromine, and iodine; or the chlorates, bromates, iodates, hypochlorites, etc. h. Compounds with the acids of sulphur, selenium, and tellurium; or the sulphates, sulphites, etc. i. Compounds with carbonic acid; or the carbonates. . j. Compounds with the acids of phosphorus, arsenic, and antimony ; or the phosphates, arsenates, etc. k. Compounds with silicic acid; or the silicates. I. Compounds with boric acid; or the borates. It is more important to become acquainted with the general methods of preparation and the general properties of the more important compounds than to learn details concerning many individual members of each class. Only those compounds which illustrate general principles, or which, owing to some application, happen to be of special interest, need be dealt with in this book. The acids of which the salts are derivatives are already known tf us, and in dealing with the acids frequent refer- ence has been made to the methods of making the salts, and to some of their more important properties. It will be well, before taking up the metals systematically, briefly to treat of the general methods of preparation, and the general properties of the different classes of metallic com- pounds. It must be borne in mind, however, that the only way to become familiar with these substances and their relations is by working with them in the laboratory. Chlorides.— The chlorides, as well as the fluorides, bromides, and iodides, may be regarded as the salts of CHLORIDES. 433 hydrochloric, hydrofluoric, hydrobromic, and hydriodic acids, or simply as compounds of the metals with the members of the chlorine family. The most important of these compounds are the chlorides, and these 'well illus- trate the conduct of the others. The chlorides are made by treating a metal with chlo- rine, or with hydrochloric acid; by treating an oxide or a hydroxide with hydrochloric acid; by treating an oxide with chlorine and a reducing agent, like carbon; by treat- ing a salt of a volatile acid with hydrochloric acid; by treating a salt of an insoluble acid with hydrochloric acid; by adding hydrochloric acid or a soluble chloride to a solu- tion containing a metal with which chlorine forms an insoluble compound; and by adding to a solution of a chloride a salt, the acid of which forms with the metal of the chloride an insoluble salt, while the metal contained in it forms with chlorine a soluble chloride. Only two of the above methods are peculiar to chlorides. These are the treatment of the metals with chlorine, and the treatment of oxides with chlorine and a reducing agent. The others involve principles which are also in- volved in the preparation of all salts, and they may there- fore be treated of in a general way. The formation of chlorides by direct treatment of the metals with chlorine is the simplest method of all. It has been illustrated in studying chlorine. It was found that chlorine combines with other elements with great ease. Thus, iron, copper, and tin combine with it, as repre- sented in the following equations : Cu +C1 =CuCl; 2Pe + 3C1 2 = 2Fe01 3 ; Sn + 2C1 2 = SnCl 4 . The preparation of chlorides by treating oxides with chlorine and a reducing agent has been illustrated in the making of boron trichloride. It is used in making aluminium trichloride, For this purpose, chlorine is 434 COLLEGE CHEMISTRY. passed over a heated mixture of aluminium oxide and charcoal, when reaction takes place according to the fol- lowing equation: A1 2 3 + 3C + 3C1 2 = 2A1C1 S + 3C0. The other methods for preparing chlorides are, as has been said, general in character and are applicable to most salts. Formation of Salts in General. 1. By treating a metal with an acid. — This is the simplest method. It has been illustrated in the prepara- tion of zinc sulphate by the action of zinc on sulphuric acid" Zn -f H a S0 4 = ZnS0 4 + H a . Other common examples are those represented in the fol- lowing equations: Fe + H 2 SO, = FeS0 4 + H a ; Zn + 2HC1 = ZnCl a + H„. 2. By treating an oxide or a hydroxide with an acid. — This is of more general application than the preceding method. As it has been studied in some detail in connec- tion with the subject of salts (see pp. 161-166), it need not be further treated of here. 3. By treating the salt of a volatile acid with another acid. — This method has been repeatedly illustrated in the decomposition of carbonates and nitrites by acids in general. While carbonic acid and nitrous acid themselves are perhaps not formed in these reactions, and we cannot say that the carbonates and nitrites are salts of volatile acids, yet the decomposition products of these acids are volatile at ordinary temperatures. The decomposition of carbonates by acids has been pretty fully studied, though attention was not directed to the fact that this kind of action may be utilized for the purpose of making salts. As some carbonates occur in large quantity in nature or in FORMATION OF SALTS IN GENERAL. 435 the market, salts are frequently made by treating them with acids. Thus, magnesium sulphate is made by treat- ing magnesium carbonate with sulphuric acid: MgC0 3 + H 2 S0 4 = MgSO, + H 2 + CO,; and calcium chloride is made by dissolving calcium car- bonate in hydrochloric acid: CaC0 3 + 2HC1 = CaCl 2 + H 2 + C0 2 ; etc. These reactions proceed to the end for the reason that one of the products is volatile and is carried out of the field of action (see p. 412). 4. By treating a salt of an insoluble acid with another acid. — This case does not occur practically, as there are no common, insoluble acids. The principle involved is illustrated to some extent by the decomposition of a soluble silicate. Sodium silicate, for example, is soluble. When its solution in water is treated with an acid the silicic acid is partly precipitated, as we have seen : ]STa 2 Si0 3 + 2HC1 + H 2 = Si(OH) 4 + 2NaCl. The silicic acid formed is, however, not perfectly in- soluble in water, so that the reaction is not complete. In any case the reaction is not one that is used for the prep- aration of salts. 5. By the action of two salts upon each other. — This method can be best described by means of an example. Suppose it is desired to prepare copper chloride by the action of two salts upon each other. Copper chloride is soluble. If copper sulphate and barium chloride are brought together in solution, the products are insoluble barium sulphate and soluble copper chloride: CuS0 4 + BaCl 2 = BaS0 4 + CuCl 2 . By simply filtering off the barium sulphate, a solution of copper chloride is obtained. 436 COLLEGE CHEMISTRY. 6. By precipitation. — This method is illustrated in the formation of barium sulphate, referred to in the last para- graph. Obviously, it is applicable only to difficultly solu- ble or insoluble salts. Many carbonates and phosphates can be made in this way. General Properties of the Chlorides Most of the chlo- rides of the metals are soluble in "water without decom- position, though many of them are decomposed when heated to a sufficiently high temperature with water. It will be remembered that the chlorides of the non-metallic or acid-forming elements are decomposed by water, yield- ing the corresponding oxides or hydroxides. The chlorides of some elements which are partly basic and partly acid are only partly decomposed. This is illustrated by the chloride of antimony, which with water forms an oxy- chloride: SbCl 3 + H 2 = SbOCl + 2HC1. The chlorides of the most strongly marked metals, like potassium, sodium, etc., are not decomposed by water. Calcium chloride dissolves with great ease, and, if the solution is evaporated, the chloride is again obtained. If, however, the attempt is made to drive off all the water by heat, some of the chloride is converted into the oxide as represented in the equation CaCl a + H 2 =- CaO + 2H01. Magnesium chloride is completely decomposed, if its solu- tion in water is evaporated to dryness, the action being the same in character as that which takes place in the case of calcium chloride. The chlorides of iron and aluminium and of many other metals act in the same way. Silver chloride and mercurous chloride, HgCl, are insoluble in water. Lead chloride is difficultly soluble in water. If, therefore, on adding hydrochloric acid or a soluble chlo- ride to a solution, a precipitate is formed, the conclusion is generally justified that one or more of the three metals THE SO-CALLED DOUBLE CHLORIDES. 437 — silver, lead, or mercury — is present. By taking into account the differences between these chlorides, it is not difficult to decide of which of them a precipitate consists. The chlorides are for the most part decomposed when treated with sulphuric acid, as has been shown in the action of sulphuric acid upon sodium chloride. Under these circumstances hydrochloric acid is given off, and the sulphate of the metal with which the chlorine was in com- bination is formed. In general, the reaction is represented by such equations as the following: 2MC1 + H 2 S0 4 = M,S0 4 + 2HC1; M"C1 S + H 2 S0 4 = M"S0 4 + 2HC1; etc. Under ordinary circumstances, chlorides are not decom- posed by any acid except sulphuric acid. The So-called Double Chlorides and Similar Compounds of Fluorine, Bromine, and Iodine. — These compounds and their relations to the oxygen salts have been repeatedly referred to. Many chlorides combine with the chlorides of the stronger metals, like sodium and potassium, form- ing well-characterized compounds. Generally, these double chlorides are analogous to the oxygen salts in com- position, differing from them only by containing two atoms of chlorine in the place of each of the oxygen atoms. As examples of these salts of the chloro-acids those which are formed by the chlorides of platinum, antimony, chromium, and gold may be mentioned. Platinic chlo- ride, PtCl 4 , combines with other chlorides, forming salts of the general composition expressed by the formula Pt01 4 + 2MC1, or M 2 PtCl 6 . Different Chlorides of the Same Metal. — Just as sul- phur, selenium, phosphorus, and the other acid-forming elements combine with chlorine and the other members of the chlorine group in more than one proportion, so many of the metals combine with the members of the chlorine group in more than one proportion. Thus, mercury forms the two chlorides, HgCl 2 and HgCl, known respectively as mercuric and mercurous chlorides; iron forms the two 43 8 COLLEGE CHEMISTRY. chlorides FeCl, and FeCl a , known as ferric and ferrous chlorides; and tin forms stannic chloride, SnCl 4 , and stannous chloride, SnCl a . The conversion of a higher chloride into a lower one is called an act of reduction. The change can generally be effected by means of nascent hydrogen : Sn01 4 + 2H = SnCl a + 2HC1. FeCl s + H = FeCl a + HC1. The conversion of a lower chloride into a higher one is generally spoken of as an act of oxidation, for the reason that it is most commonly effected by the action of oxygen. Thus the most convenient way to transform ferrous chlo- ride into ferric chloride is to treat it in solution in hydro- chloric acid with an oxidizing agent, when a double action takes place, as represented in the following equation: 2FeCl a + 2HC1 + = 2FeCl 8 + H a O. The same change can be effected by the direct action of chlorine : FeCl, + CI = FeCl,. In this case it would obviously be incorrect to speak of the process as one of oxidation. Another method of reduction, besides that referred to above, involving the action of nascent hydrogen, is that illustrated in the equation 2HgCl i) + SnCl a = 2HgCl + SnCl 4 . In this case mercuric chloride is changed to mercurous chloride by the action of stannous chloride. The latter unites with chlorine so readily that it extracts it from some other chlorides, and is itself transformed into stannic chloride. While, therefore, we say that the stannous chloride reduces the mercuric chloride, it is equally true to say that the mercuric chloride chlorinates the stannous chloride. OXIDES. 439 Oxides. — The oxides occur extensively in nature, and are among the most common ores of some of the impor- tant metals. The oxides of iron, tin, and manganese, for example, occur in nature. They can be made by oxidiz- ing the metals, by heating nitrates, carbonates, and hydroxides, and by heating some sulphides in contact with the air. When magnesium is burned it is converted into mag- nesium oxide: Mg + = MgO. When lead nitrate is heated it gives off oxygen and an oxide of nitrogen, and lead oxide is left behind : Pb(N0 3 ) 2 = PbO + 2N0 2 + 0. When calcium carbonate is heated it yields calcium oxide and carbon dioxide : CaC0 3 = OaO + CO, When aluminium hydroxide, A1(0H) 3 , is heated it loses water, and aluminium oxide is left behind : 2A1(0H) S = ALO, + 3H a 0. The sulphide of iron, when heated in contact with the air, or "roasted," is converted into ferric oxide and sulphur dioxide. Most of the oxides of the metals are insoluble in water. Those of Group A, Family I, are soluble, but are con- verted by water into the corresponding hydroxides. The oxides are acted upon generally by acids forming the corresponding salts. If the salt with a certain acid is insoluble, the salt is not formed by the action of that acid on the oxide unless the acid or its anhydride is fusible and not volatile, when by fusing them together the salt is i formed. ^\\ Different Oxides of the Same Metal. — Just as there are i different chlorides of the same metal, so there are different 440 COLLEGE CHEMISTRY. oxides, and indeed there is greater variety among these than among the chlorides. Iron forms three oxides, ferric oxide, Fe 2 3 , ferroso-ferric oxide, Fe 3 4 , and ferrous oxide, FeO ; mercury forms the two oxides HgO and Hg a O ; etc. The lower oxides are converted into the higher by oxidation, and the higher into the lower by reduction. The higher oxides of several of the metals are acidic. This is markedly so in the case x>i chromium and man- ganese. • Hydroxides. — The hydroxides are formed by treating oxides with water and by decomposing salts by adding soluble hydroxides to their solutions. In general, when- ever a salt is decomposed by a strong base, the base of the salt separates in the form of the hydroxide. The forma- tion of a hydroxide by the action of water on an oxide is well illustrated by the action of water on lime or calcium oxide, a process familiarly known as slaking : CaO + H 2 = Ca(OH) 3 . Most of the hydroxides of the metals are insoluble in water. If a soluble hydroxide is added to a solution con- taining a metal whose hydroxide is insoluble, the latter is precipitated. Thus, if a solution of sodium hydroxide is added to a solution of a magnesium salt, magnesium hydroxide is precipitated: MgS0 4 + 2NaOH = Na 2 SO, + Mg(OH) a . So, also, when a solution of a ferric salt is treated with sodium hydroxide, a precipitate of ferric hydroxide is formed : PeCl 3 + 3NaOH = 3NaCl + Fe(OH),. Only the hydroxides of the members of the potassium family, and some of the members of the calcium family, are soluble in water. The hydroxides of sodium and potas- sium are called alkalies. The solution of ammonia in water acts like a soluble hydroxide, and probably contains DECOMPOSITION OF SALTS BY BASES. 441 ammonium hydroxide, NH 4 (OH), formed by the action of water on ammonia: NH 5 + H 2 = NH 4 (OH). Now, when any one of the soluble hydroxides is added to a salt containing any metal that does not belong to the potassium or calcium family, an insoluble compound is formed. Decomposition of Salts by Bases The decomposition of salts by bases is analogous to the decomposition by acids. When a soluble base acts upon a salt, there are four possible kinds of action : 1. The base from which the salt is derived may be vola- tile, or may break up, yielding a volatile product. In this case, decomposition takes place and the volatile base is given off. This is not a common case except among the compounds of carbon. The one illustration which we have had is the decomposition of ammonium salts by calcium hydroxide and sodium hydroxide, when the vola- tile compound ammonia, NH 3 , is given off. 2. The hydroxide, or base from which the salt is derived, may be insoluble or difficultly soluble in water, and not volatile. In this case, if both the salt and the base are in solution, decomposition takes place, and the insoluble or difficultly soluble hydroxide, or base, is precipitated. This has already been illustrated. 3. The base from which the salt is derived may be solu- ble and not volatile. This is the case, for example, when sodium hydroxide is added to a solution of potassium nitrate. Here sodium nitrate, potassium nitrate, sodium hydroxide, and potas- sium hydroxide may all be present in the solution, and investigation has shown that all are present and that the quantity of each depends upon the masses of the sub- stances brought together, and upon their affinities. 4. The fourth case is that in which a soluble hydroxide 442 COLLEGE CHEMISTRY. forms an insoluble salt with the acid of a soluble salt, leaving a soluble hydroxide in solution. This is illustrated by the action of calcium hydroxide on a solution of sodium carbonate, when insoluble calcium carbonate is thrown down, and sodium hydroxide remains in solution, as represented in the equation Na a COj + Ca(OH) 2 = 2NaOH + CaCO s . Some basic hydroxides, which are precipitated by solu- ble hydroxides, have a weak acid character, and, after they are precipitated, they redissolve in an excess of the soluble hydroxide. This is true, for example, of the hydroxides of aluminium, chromium, and lead. The salt-like com- pounds thus formed are generally quite unstable. The precipitation and subsequent solution of the hydroxides of the three metals named take place thus : A1C1, + 3NaOH = Al(OH), + 3NaCl; Al(OH), + 3NaOH = Al(ONa) s + 2H 2 0; CrCl, + 3NaOH = Cr(OH) 3 + 3NaCl; Cr(OH) 3 + 3NaOH = Cr(ONa), + 3H 2 0; Pb(N0 2 ) 2 + SNaOH = Pb(OH) 2 + 2NaNO s ; Pb(OH) 2 + 2NaOH = Pb(ONa) 2 + 2H 2 0. In some cases where a soluble hydroxide is added to a salt, an oxide is precipitated instead of the hydroxide. This is analogous to the formation of an anhydride of an acid instead of the acid itself, as when carbonates, sul- phites, and nitrites are decomposed. When a silver salt is treated with a soluble hydroxide, silver oxide is at once precipitated. The same is true of mercury salts : SAgNO, + 2KOH = Ag 2 -f H 2 + 2KN0 3 ; HgCl 2 + 2NaOH = HgO + H 2 + SNaCl. It is probable that the first product is the hydroxide, and that this breaks down into the oxide and water: SULPHIDES. 443 2AgN0 3 + 2K0H = 2AgOH + 2KN0 8 ; 2AgOII = Ag 2 + H 2 0; HgCl 2 + 2NaOH = Hg(OH) 2 + 2NaCl; Hg(OH) 2 = HgO + H 2 0. Some hydroxides are converted into the oxides by simply boiling the liquids in which they are suspended. Thus, when a salt of copper is treated with a soluble hydroxide, copper hydroxide is first precipitated; but if the solution in -which it is suspended is boiled, it is soon changed to the oxide: CuS0 4 + 2NaOH = Na 2 S0 4 + Cu(OH) 2 ; Cu(OH) 2 = 2 CuO + H 2 0. The hydroxides corresponding to some of the higher oxides of the metals, as those of chromium and manganese, are acids. The hydroxides of most of the metals are decomposed by heat into water and the corresponding oxides. Those of the alkali metals, as potassium and sodium, are, how- ever, not decomposed by heat. Sulphides. — Many sulphides are found in nature as, for example, iron pyrites, FeS 2 ; lead sulphide, or galenite, PbS; copper pyrites, FeCuS 2 ; etc. They are made in the laboratory by heating metals with sulphur; by treating solutions of salts with hydrogen sulphide; by treating solutions of salts with soluble sulphides; and by reducing sulphates. Attention has been called to the fact that the sulphides are analogous in composition to the oxides, and that they are to be regarded as salts of hydrogen sulphide formed by substituting metals for the hydrogen of the acid. The formation of sulphides by the direct combination of sulphur with the metals is shown in the formation of lead sulphide and copper sulphide : Pb +S = PbS; 2Cu + S = C^S. 444 COLLEGE CHEMISTRY. The formation of sulphides by the action of hydrogen sulphide upon solutions of salts was discussed at some length under Hydrogen Sulphide (which see). The ex- tensive use made of this reaction in chemical analysis was also referred to. The action of soluble sulphides or solutions of salts is in general the same as that of hydrogen sulphide, but in some cases, in which the former will not act, the latter will. Thus, hydrogen sulphide will not precipitate iron sulphide from a solution of an iron salt, because iron sul- phide is easily acted upon by dilute acids. When, for example, hydrogen sulphide is passed into a solution of ferrous chloride, it naturally tends to form the sulphide FeS: FeCl a + H a S = FeS + 2HC1. But ferrous sulphide, FeS, is acted upon by dilute hydro- chloric acid, and is converted by it into ferrous chloride and hydrogen sulphide: FeS + 2HC1= FeCl a + H a S. It is therefore obvious that the first reaction cannot take place. If, however, a soluble sulphide, as sodium or ammonium sulphide, is added to a solution of an iron salt, iron sul- phide is precipitated, as in this case no free acid is formed. Thus, when ferrous chloride and ammonium sulphide are brought together the reaction takes place as represented in the equation FeCl a + (NH 4 ) a S = FeS + 2NH 4 C1. In ammonium chloride the ferrous sulphide is no* soluble. The formation of a sulphide by reduction of a sulphate is illustrated by the formation of barium sulphide by heat- ing a mixture of barium sulphate and charcoal : BaSO, + 40 = BaS + 400. HYDROSULPHIDES. 445 The sulphides of the alkali metals are soluble in water. Those of the other metals are insoluble. It should be remarked, however, that aluminium and chromium do not form sulphides, or, at least, if they do, the compounds are decomposed by water into hydroxides and hydrogen sul- phide. Barium sulphide is decomposed by water, and probably magnesium sulphide also. The sulphides are stable when heated without access of air; but if heated in the air they are converted into oxides of the metals and sulphur dioxide, or, in some cases, they take up oxygen and are converted into sulphates. The conversion of sulphides into oxides and sulphur dioxide by heating them in contact with the air has been repeatedly referred to. The process is carried on on the large scale in the preparation of iron ores for reduction, and is called roasting. The conversion of a sulphide into a sulphate by heating is a simple process of oxidation. Copper sul- phide is converted into the sulphate when heated for some time: CuS + 40 = OuS0 4 . This is the reverse of the reaction mentioned by which a sulphate is converted into a sulphide by reduction. Some sulphides, as those of sodium, potassium, and ammonium, take up sulphur in much the same way that they take up oxygen, and form the polysulphides. The two reactions appear to be analogous : K 6 S + 40 = K 2 S0 4 ; K S S + 4S = K a SS„ or K 8 S S . Hydrosulphides. — The hydrosulphides bear the same relation to the sulphides that the hydroxides bear to the oxides. They are not, however, as numerous nor as easily obtained as the hydroxides. The only hydrosulphides known are derived from the members of the potassium and calcium groups, and these are soluble. They are formed by saturating solutions of 446 COLLEGE CHEMISTRY. the corresponding hydroxides with hydrogen sulphide. Potassium hydrosulphide is formed thus : KOH + H 2 S = KSH + H a O. Ammonium hydrosulphide is formed thus : NH.OH + H 2 S = NHJ3H + H a O. It also appears probable that whenever a sulphide dissolves in water it is converted into a hydrosulphide and a hy- droxide. Thus it semes to be true that potassium sulphide is converted into the hydrosulphide and hydroxide: K a S + H 2 = KSH + KOH. Sulpho-salts. — The relation of the sulpho-salts to the sulphides has already been explained. It is like that of the ordinary oxygen salts to the oxides, and that of the chloro-salts, or double chlorides, to the chlorides. They are formed by dissolving the sulphides of certain metals, particularly tin, arsenic, and antimony, in the sulphides of the members of the potassium group : As a S s + 3K 2 S = 2K 3 AsS,; As 2 S 6 + 3K 2 S = 2K 3 AsS 4 ; SnS 2 + K 2 S = K 2 SnS 3 , etc. These sulpho-salts are decomposed by the ordinary acids, the insoluble sulphides being precipitated thus: 2K 3 AsS 3 + 6HC1 = As 2 S 3 + 6KC1 + 3H 2 S- 2K 3 AsS 4 + 6HC1 = As 2 S 5 + 6KC1 + 3H 2 S. Nitrates. — The nitrates are formed by dissolving the metals in nitric acid, and by treating oxides, hydroxides, carbonates, and some other easily decomposed salts with nitric acid. The action of nitric acid upon metals was discussed under the head of Nitric Acid (which see). It was pointed out that the acid gives up a part of its oxygen CHLORATES. 447 to the metal and forms different oxides, according to the conditions. Thus, when the acid acts upon copper the main product of the reduction is nitric oxide, but by changing the concentration of the acid a considerable quantity of nitrous oxide is formed. When zinc is dis- solved in nitric acid a part of the acid is reduced to ammonia. The nitrates are soluble in water, and all are decom- posed by heat. Some of them when heated lose only a third of their oxygen and are reduced to nitrites. This is true of potassium nitrate, the decomposition of which is represented by the equation KNO s = KNO, + 0. Most of the nitrates, however, are decomposed further, forming oxides. This has been shown in the case of lead nitrate, which when heated is converted into lead oxide, while nitrogen peroxide and oxygen are given off: Pb(NO s ) a = PbO + 2N0, + 0. If the oxide of the metal is decomposed by heat, as in the case of mercury, of course the product will be the metal. Chlorates. — These salts, except potassium chlorate, are not commonly met with. Potassium chlorate is manufac- tured in large quantity, and the other chlorates are generally made from it. The chlorates are soluble in water, and are decomposed by heat more easily than the nitrates are. They are first converted into perchlorates, and these are further decomposed by higher heat into chlorides and oxygen. The hypochlorites are formed by treating some of the metallic hydroxides in dilute solution with chlorine. This has been illustrated in the formation of "bleaching powder," which contains calcium hypochlorite or a com- pound closely related to it. The hypochlorites, like the chlorates, are decomposed by heat. 448 COLLEGE CHEMISTRY. Sulphates. — The general relations of the sulphates to sulphuric acid were treated of under Sulphuric Acid (which see). Some of these salts occur in nature in large quantity, as those of calcium and barium. The former is known as gypsum, the latter as heavy spar. Sulphates are made by treating metals, metallic hydroxides or oxides, carbonates, etc., with sulphuric acid; and by treating a solution containing a metal whose sulphate is insoluble, with sulphuric acid or a soluble sulphate. Zinc and iron give hydrogen and a sulphate when treated with sulphuric acid: Zn + H 2 S0 4 = ZnS0 4 + H,; Fe + H 2 S0 4 = FeS0 4 + H 2 . This kind of action takes place whenever a metal is dis- solved in sulphuric acid at the ordinary temperature. If, however, the temperature is raised the displaced hydrogen acts upon some of the sulphuric acid, or the metal extracts some of the oxygen of the acid, reducing it partly to sul- phurous acid, when sulphur dioxide is given off. This happens in the case of copper, as has been pointed out. It may be represented either by these equations: Cu + H 2 S0 4 = CuS0 4 + H 2 ; H 2 S0 4 + 2H = SO a + 2H 2 0; or by these : Cu + H 2 S0 4 = CuO + S0 2 + H,0; CuO + H 2 S0 4 = CuS0 4 + H 2 0. The action of sulphuric acid on metallic hydroxides has been fully described. Most sulphates are soluble in water. The sulphates of barium, strontium, and lead are insoluble in water, and the sulphate of calcium is difficultly soluble. Therefore, if sulphuric acid or a soluble sulphate is added to a solu- tion containing either of the metals, barium, strontium, or lead, a precipitate is formed. A precipitate is also CARBONATES. 449 formed when a concentrated solution o^ a calcium salt is treated in the same way. When heated with charcoal in the reducing flame of the blowpipe, sulphates are reduced to sulphides : K a S0 4 + 4C = K a S + 4CO, or K a S0 4 + 2C = K 3 S + 2CO a . Sulphites are made from sodium or potassium sulphite, which are made by treating sodium or potassium hydroxide in solution with sulphur dioxide : 2NaOH + SO a = Na,SO, + H a O. All sulphites are decomposed by the common acids, sul- phur dioxide being given off: Na a SO, + H a SO, = Na a S0 4 + H a O + SO a . The sulphites are changed to sulphates by oxidation. Thus, sodium sulphite is changed to the sulphate when its solution is allowed to stand in contact with the air: Na a SO s + = Na a S0 4 . The sulphites, like the sulphates, are reduced to sul- phides. Carbonates. — Many carbonates are found in nature, some of them in great abundance and widely distributed. The principal one is calcium carbonate. They are made by passing carbon dioxide into solutions of hydroxides, and by adding soluble carbonates to solutions of salts con- taining metals whose carbonates are insoluble. The formation of carbonates by the action of carbon dioxide on a solution of a hydroxide is illustrated in the case of potassium hydroxide : 2KOH + C0 3 = K a C0 8 + H a O. 45° COLLEGE CHEMISTRY. The formation of calcium carbonate takes place in the same way, but the carbonate formed is insoluble : Ca(OH) 2 + C0 3 = CaC0 3 + H 2 0. If in either case the action is continued, the norma carbonate first formed is converted into acid carbonate : K 2 CO s + CO, + H 2 = 2KHCO s ; CaCO s +CO a + H 2 = Ca| (00a)a! All normal carbonates except those of the members of the potassium family are insoluble, and are decomposed by heat into carbon dioxide and the oxide of the metal. The decomposition of calcium carbonate into lime and carbon dioxide is an example : CaC0 3 = CaO + CO,. When a soluble carbonate is added to a solution of a calcium, barium, or strontium salt, the corresponding in- soluble carbonates are precipitated. When a magnesium salt is treated with a soluble carbonate, however, a basic carbonate is precipitated. Phosphates. — Calcium phosphate is very abundant in nature, and a few other phosphates are also found. The methods used for making phosphates are the same as those used in making salts in general. The normal phosphates of all the metals except the members of the potassium family are insoluble in water. The normal phosphates, as a rule, are not changed by heat. The secondary phosphates, such as secondary sodium phosphate, HNa 2 P0 4 , lose water when heated, and yield pyrophosphates: 2HNa 2 P0 4 = Na 4 P 2 0, + H 2 0. Sodium pyrophosphate. SILICATES. 451 Those phosphates in which only one-third of the hydrogen is replaced by metal — as, for example, primary sodium phosphate, H a NaP0 4 — lose water when heated, and yield metaphosphates: H.NaPO, = NaPO, + H a O. Sodium metaphospkate. Neither the pyrophosphates nor the metaphosphates are changed by heat. Silicates. — The silicates, as has been stated, are very widely distributed in nature. Those which are most abundant are the feldspars and their decomposition- products. The principal feldspar is a complex silicate of aluminium and potassium, of the formula KAlSi 3 8 , derived from the polysilicic acid H 4 Si 3 B , which is formed from three molecules of normal silicic acid by the loss of four molecules of water: 3Si(OH) 4 = H 4 Si 3 8 + 4H 2 0. Silicates can be made by heating together, at a high temperature, silicon dioxide, in the form of sand, and basic oxides or carbonates : CaO + SiO a = CaSiO s ; Na 3 C0 3 + Si0 3 = Na 2 SiO s + CO,. Only the silicates of the members of the potassium group are soluble in water. When these are treated in solution with dilute acids, they are decomposed, as has been explained under Silicic Acid (which see). Some silicates, which are insoluble in water, are decom- posed by the ordinary acids, such as sulphuric and hydro- chloric acids, the silicic acid separating as a difficultly soluble substance, which, if dried on the water-bath, becomes insoluble. Many silicates, which are not acted upon by strong acids, are decomposed by fusion with sodium or potassium 452 COLLEGE CHEMISTRY. carbonate, when the silicate of potassium or sodium and the oxide of the metal contained in the silicate are formed. Silicates which are not decomposed in either of the ways mentioned yield to hydrofluoric acid. The action con- sists in the formation of the gas, silicon tetrafluoride, Si.P 4 , and the fluorides of the metals present. Thus, the reaction in the case of feldspar takes place in accordance with the equation, KAlSi 3 8 + 16HP = KP + A1F ? + 3SiP 4 + 8H„0. The silicon fluoride is given off as a gas, and the fluorides formed are soluble in water. Hence, hydrofluoric acid is said to dissolve the silicates. EXPERIMENTS. Chlorides, Bromides, and Iodides. ^Experiment 201.— Dissolve a small crystal of silver nitrate in pure water. Add to a small quantity of this solution in a test- tube a few drops of dilute hydrochloric acid. The white sub- stance thus precipitated is silver chloride, AgCl. To another small portion of the solution add a few drops of a dilute solution of common salt, or sodium chloride, NaCl. The white substance produced in this case is also silver chloride. Add ammonia to each tube. If sufficient is added the precipitates will dissolve. On adding enough hydrochloric acid to these solutions to com- bine with all the ammonia the silver chloride is again thrown down. On standing exposed to the light both precipitates change color, becoming finally dark violet. The reactions involved in the above experiments are these : In the first place, when hydro- chloric acid is added to silver nitrate this reaction takes place : + - + - + Ag + NO, + H + CI = AgCl + H + NO,. When sodium chloride is added this reaction takes place : Ae + NO, + Na + CI = AgCl + Na + NO,. EXPERIMENTS: HYDROXIDES. 453 In the first reaction the ions of nitric acid are formed ; in the second, the sodium and silver exchange places. In addition to the insoluble silver chloride, there are formed at the same time the ions of the soluble salt, sodium nitrate. On adding ammonia the silver chloride forms with it a compound which is soluble in water ; and on adding an acid, the ammonia combines with it, leaving the silver chloride uncombined and therefore insoluble. Extensive use is made of insoluble compounds for the purpose of detecting substances in analysis. The only insoluble chlorides are those of silver, lead, and mercury.* If, therefore, on adding hydrochloric acid or a soluble chloride to a solution, a precipitate is formed, the conclusion is justified that one or more of the three metals — silver, lead, or mercury — is present. By taking account of the- differences in the properties of these chlorides it is not dif- ficult to decide of which of them a precipitate consists. Hydroxides. Experiment 202. — To some pieces of freshly-burnt lime add enough cold water to cover it. The action which takes place is represented by the equation CaO + H 3 = CaCOH),. The process is known as slaking. i\ Experiment 203. — To a small quantity of a dilute solution of magnesium sulphate add a dilute solution of caustic soda. The white precipitate is magnesium hydroxide. [Would you expect this precipitate to be soluble in sulphuric acid ? in hydrochloric acid ? in nitric acid ?] The answers follow from these considera- tions : When acids act upon hydroxides, salts are formed ; mag- nesium sulphate is soluble, as is seen by the fact that we started with a solution of this salt ; the only insoluble chlorides are those of silver, lead, and mercury ; all nitrates are soluble. When a solution of an iron salt is treated with sodium hydrox- ide a precipitate of iron hydroxide is formed : FeCl, + 3NaOH = EeO s H 3 + 3Na01. \» Experiment 204. — To a dilute solution of that chloride of iron which is known as ferric chloride add caustic soda. The reddish * There are two chlorides of mercury. Only one of them, mer- curous chloride, is insoluble. 454 COLLEGE CHEMISTRY. precipitate which is formed is ferric hydroxide. [From the gen- eral statements made above, would you expect this precipitate to be soluble in hydrochloric acid ? in nitric acid ? Try each. Is it soluble in sulphuric acid ?] ^ Experiment 205. — Add to a solution of an aluminium salt sodium hydroxide. After a precipitate is formed continue to add the sodium hydroxide. Perform similar experiments with a chromium and with a lead salt. Boil each of the solutions ob- tained. Treat a solution of copper sulphate with sodium hydrox- ide in the cold. Heat. Sulphates. ■^ Experiment 206. — Make a dilute solution of barium chloride, of lead nitrate, of strontium nitrate. To a small quantity of each —in a test-tube add a little sulphuric acid. [What remains in solu- tion ?] Make a somewhat concentrated solution of calcium chlo^ ride. To this add sulphuric acid. ["What is in solution ?] Add more water, and see whether the precipitate will dissolve. The formulas of the salts used in the experiments are barium chloride, BaOU ; lead nitrate, Pb(NO a )j ; strontium nitrate, Sr(N0 3 ) a . [Write the equations expressing the reactions.] If to the solutions of the salts any soluble sulphate is added instead of sulphuric acid, the same insoluble sulphates will be formed. The sulphates of iron, copper, sodium, and potassium are among the soluble sulphates. Make dilute solutions of small quantities of each of these, and add them successively to the solutions of barium chloride, lead nitrate, and strontium nitrate. The formula of iron sulphate is FeSO» ; of copper sulphate, CuSOi ; of sodium sulphate, Na a SO« ; and of po- tassium sulphate, K 2 S0 4 . Write the equations representing the reactions which take place in the above experiments. It need hardly be explained that the action consists in an exchange of places on the part of the metals. Thus, when the soluble salt iron sulphate, FeSO< , is brought together with the soluble salt barium chloride, Ba01 3 , the insoluble salt barium sulphate, BaS0 4 r and the soluble salt iron chloride, FeCl a , are formed : FeSO. + Bad, = FeCl„ + BaS0 4 . Reduction of Sulphates to Sulphides" -^Experiment 207. — Mix and moisten a little sodium sulphate and finely-powdered charcoal. Heat the mixture for some time in the reducing flame. After cooling scrape pf£ the salt, dissolve EXPERIMENTS : CARBONATES. 455 it in a few cubic centimetres of water, and filter through a small filter. If the change to the sulphide has taken place, sodium sul- phide, Na a S, is in solution. A solution of a sulphide when added to a solution containing copper gives a black precipitate of copper sulphide. Try this ; also try the action on the solution ' of the salt of copper of some of the sulphate - from which the sulphide was made. V Cakbonates. Experiment 208. — The formation of carbonates by the addi- tion of soluble carbonates to solutions of salts of metals whose carbonates are insoluble, is illustrated by the following experi- ments : Make solutions of copper sulphate, iron sulphate, lead nitrate, silver nitrate, calcium chloride, barium chloride. Add to each a little of a solution of a soluble carbonate, as sodium carbonate, potassium carbonate, ammonium carbonate. Note the result in each case. Filter off all the precipitates and prove that they are carbonates. This may be done by treating them with dilute acids, which decompose them, causing an evolution of carbon dioxide, which can be detected by passing a little of it into lime-water. In some of the cases mentioned the insoluble salts formed are basic carbonates, as, for example, those of cop- per and magnesium. The salts of silver, calcium, and barium are the normal carbonates Ag a C0 3 , BaCOa , and CaCOs . CHAPTER XXV. ELEMENTS OF FAMILY I, GROUP A: THE ALKALI METALS :— LITHIUM.— SODIUM.— POTAS- SIUM.— RUBIDIUM.— CAESIUM.— AMMONIUM. General. — The elements of this group which are most abundant in nature are sodium and potassium. While lithium occurs in considerable quantity, the two remaining elements, rubidium and caesium, have been found in only very small quantities. They are all strongly basic, their hydroxides being the strongest bases known. They form well-characterized salts with all acids, and as a rule their salts are very stable. In all their compounds they act as univalent elements, except in those which they form with hydrogen, and in their peroxides; in the latter they appear to be bivalent. Leaving these compounds out of consid- eration the general formulas of some of the other principal compounds are as follows : MOl, M,0, M 2 S, M(OH), M(SH), MN0 S , M 2 S0 4 , etc. The valence of the members of the group towards other elements is, in general, constant. The relations between the atomic weights are interest- ing. That of sodium, 23.05, is very nearly half the sum of those of lithium, 7.03, and potassium, 38.15. We have 7.03 + 39.15 „„ „„ ~ = 23.09. So, also, that of rubidium, 85.4, is approximately half the sum of those of potassium, 39.15, and csesium, 133. 39.15 + 133 ~ = 86.08. 456 POTASSIUM. 457 The specific gravity of these elements increases with the atomic weight; and their melting-points become lower as the atomic weights become higher. At. Wt. Sp. Grav. M. P. Lithium 7.03 0.594 180°. Sodium 23. 05 0. 972 95.6° Potassium 39.15 0.865 62.5° Pubidium 85.4 1.52 38.5° Caesium 133 ? ? The regularity is complete in the case of the melting- points, but as regards the specific gravities sodium is an exception to the rule. As sodium and potassium and their compounds are much more commonly met with than the other members of the group, these will form the chief subject of this chapter. Potassium, K (At. Wt. 39.15). Occurrence. — Potassium is a constituent of many min- erals, particularly of feldspar, the common variety of which, as has already been explained, is a complex silicate of aluminium and potassium. It is found also in com- /bination with chlorine as carnallite and sylvite; with sul- phuric acid and aluminium, as alum; with nitric acid, as saltpetre or potassium nitrate; and in other forms. The natural decomposition of minerals containing potassium gives rise to the presence of this metal in various forms of combination everywhere in the soil. It is taken up by the plants; and, when vegetable material is burned, the potas- sium remains behind, chiefly as potassium carbonate. When wood-ash is treated with water the potassium car- bonate is dissolved, and it can be obtained in an impure state by evaporating the solution. The substance thus obtained is called potash. In the juice of the grape there is contained a salt of potassium, mono-potassium tartrate, which is deposited in large quantity from wine. This is commonly called "crude tartar." 45 8 COLLEGE CHEMISTRY. Preparation. — Potassium was first prepared by Davy in the year 1807, by the action of a powerful electric current on potassium hydroxide. It was then for many years pre- pared by heating to a high temperature a mixture of potassium carbonate and carbon, obtained by heating in a closed vessel ordinary mono-potassium tartrate deposited from wine. K a C0 3 + 2C = 2K + 3CO. The metal is now prepared by electrolysis of potassium hydroxide or potassium chloride. Properties. — Potassium is a light substance, which floats on water. Its freshly cut surface has a bright metallic lustre, almost white; it acts energetically upon water, causing the evolution of hydrogen, which, together with some of the potassium, burns, while potassium hydroxide is formed at the same time. This reaction has been studied in connection with hydrogen. In consequence of its action upon water, potassium cannot be kept in the air. It is kept under some oil, as petroleum, upon which it does not act. In an atmosphere upon which it does not act, as, for example, hydrogen, it can be distilled. Its vapor is green. Its specific gravity is 0. 865 ; its melting- point 62.5°. It combines with chlorine and bromine with great energy, and has the power to extract chlorine from its compounds. It can, therefore, be used for the purpose of isolating some elements, as, for example, magnesium and aluminium, whose oxygen compounds cannot be reduced by the ordinary methods. As, however, sodium is generally used for this purpose instead of potassium, on account of its lower price, the action will be referred to more at length under Sodium. Although the metal is converted into vapor, no reliable determination of the specific gravity of the vapor has been made, for the reason that the vessels which have been used for the purpose have always been acted upon, and the results thus vitiated. Potassium Hydride, K 2 H. — This compound is formed by heating potassium in an atmosphere of hydrogen at Potassium POTASSIUM SALTS. 459 about 300°. It is a silver-white mass with a metallic lustre. It takes fire in the air. When heated, it begins to dissociate at 200°. Fluoride, KP Chloride, KC1 r., ., ,, ,, , Bromide^ KBr'" * these salts the onl y Iodide, KI one that occurs in nature in quantity is the chloride. This is found in the great salt deposits at Stassfurt, Germany, and in some other localities in the form of the mineral sylvite, which is more or less impure potassium chloride. It is also found in the form of a compound containing magnesium, potassium, and chlorine, of the formula MgCl 2 .KCl + 6H 2 0, or KMgCl, + 6H 2 0, known as camallite. The other salts of the group are made by the general methods for making salts, that is, by neutralizing the acids with the hydroxide or carbonate of potassium. It is, however, easier to make the iodide by other methods, and as there is a large demand for this salt for use in medicine and in the art of photography, several methods have been devised for its preparation. Of these, only one need be described. This consists in treating iron filings under water with iodine. Both the iron and iodine dissolve, forming ferrous iodide, Felj. If to the solution of this compound half as much iodine is added as has already been used in its preparation, ferroso-ferric iodide, Fe^, is formed and remains in solution. By adding a solution of potassium carbonate to this, reaction takes place as repre- sented in the equation: Fe s I 8 + 4K,C0 3 + 4H 2 = SKI + Pe 3 (OH) 8 + 4C0 2 . The hydroxide of iron is insoluble, and can be removed by filtration. The fact that the specific gravity of hydrofluoric acid at a low temperature corresponds to the formula H,F 2 makes it not improbable that potassium fluoride has the formula K a F 2 . This appears still more probable from the fact that 4<>° COLLEGE CHEMISTRY. there is an acid potassium fluoride of the formula KHF , or KF -\- HF. Similar acid salts have not been obtained from the other acids of the group. Properties. — All these salts are soluble, and crystallize well in cubes. The fluoride is the most easily soluble in water. If deposited from a water solution at the ordinary temperature the crystals contain two molecules of water of crystallization, and are deliquescent. The iodide is solu- ble in 0.7 part of water at the ordinary temperature, and is also soluble in alcohol (40 parts). The bromide requires about 1£ parts of water for solution at the ordinary tem- perature, and is but slightly soluble in alcohol. The chloride is soluble in 3 parts of water at the ordinary tem- perature, and is insoluble in alcohol. All are decom- posed by sulphuric acid. Applications. — Potassium chloride is extensively used for the purpose of making other potassium salts, as, for example, the nitrate and carbonate; the bromide is used in medicine ; the iodide, as stated above, is used in medi- cine and in photography. Potassium Hydroxide, KOH. — This well-known sub- stance, commonly called caustic potash, is prepared by treating potassium carbonate in solution with calcium hydroxide in a silver or iron vessel. The reaction is based upon the fact that calcium carbonate is insoluble, and that potassium carbonate and calcium hydroxide are soluble: K a CO s + Ca(OH) a = 2KOH + CaCO s . Potassium hydroxide is also prepared on the large scale by the action of the electric current on a solution of potassium chloride: 2KC1 + 2H a O = 2KOH + H, + Cl a . It is a white brittle substance. In contact with the air it deliquesces, and absorbs carbon dioxide, being com- pletely transformed into potassium carbonate. It is the strongest of the bases. It decomposes the salts of all other POTASSIUM OXIDE. 461 bases, even of those which, like sodium and lithium hydroxides, are soluble in water. Animal substances like the skin are disintegrated by the hydroxide. It has a caustic action. Instead of potassium hydroxide, the corresponding sodium compound is used wherever this is possible, as the latter is cheaper. The chief application of the potassium compound outside of the laboratory is for making soft- soap. For this purpose fats are boiled with a solution of potassium hydroxide or carbonate. Z" 1 Potassium Oxide, K 2 0.— This compound can be made by burning potassium in the air, and heating the residue to a high temperature. It is also formed by melting potassium hydroxide and metallic potassium together : 2K + 2KOH - 2K 2 + H 2 . With water it forms the hydroxide, with a marked evolu- tion of heat : K 2 + H 2 = 2KOH. Potassium also forms other oxides of which the peroxide of the formula K 2 4 is the best studied. This peroxide is the final product of the combustion of potassium in the air or in oxygen. At a high temperature it breaks down into potassium oxide, K a O, and oxygen. It also gives up its oxygen very readily to substances which are capable of oxidation, acting so energetically upon some as to cause evolution of light. Potassium Hydrosulphide, KSH, is analogous to potas- sium hydroxide. Just as the latter is made by the action of potassium on water, so the former can be made by the action of potassium on hydrogen sulphide : ' K 8 + 2H 2 S = 2KSH + H,. It is, however, obtained most readily by the action of hydrogen sulphide on a solution of potassium hydroxide: KOH + H 2 S = KSH + H a O. 462 COLLEGE CHEMISTRY. Potassium Nitrate, KN0 3 . — This salt is commonly called saltpetre. Its occurrence in nature has already been referred to under Nitric Acid (which see). When refuse animal matter is left to undergo decomposition in the presence of bases, nitrates are always the end-products. They are consequently found very widely distributed in the soil. In the East Indies the potassium nitrate formed in the neighborhood of dwellings and stables is collected, and sent into the market. The process of nitrification is carried on artificially on the large scale in the so-called "saltpetre plantations." In these, refuse animal matter is mixed with earthy material, wood ashes, etc., and piled up. These piles are moistened with the liquid products from stables. After the action has continued for two or three years the outer crust is taken off, and extracted with water. The solution thus obtained contains, besides potas- sium nitrate, calcium and magnesium nitrates. It is treated with a water-extract of wood ashes or with potas- sium carbonate, by which the calcium and magnesium are thrown down as carbonates. Much of the saltpetre in the market is made from Chili saltpetre, or sodium nitrate, by treating it with potassium chloride, advantage being taken of the fact that sodium chloride is less soluble in hot water than potassium nitrate. Molecular weights of sodium nitrate and potassium chloride are dissolved in water and the solution heated. When the solution is sufficiently concentrated, sodium chloride is deposited : NaNO a + KC1 = KNO, + NaCl. This being removed, potassium nitrate crystallizes from the mother-liquors. Potassium nitrate crystallizes in long rhombic prisms, of a salty taste. Under some circumstances it crystallizes in rhombohedro ns. j / jyhen dissolved in water it causes a /"lowering "of" thetemperature. At ordinary temperatures 100 parts of water dissolve from 20 to 30 parts of the salt; I at 100°, 100 parts dissolve 247 parts. GUWOWDER. 463 Applications. — Potassium nitrate is used as an oxidizing agent in the laboratory, and in the manufacture of fire- works. Its chief use, however, is in the manufacture of gunpowder. Gunpowder. — The value of gunpowder is due to the fact that it explodes readily, the explosion being a chemical change accompanied by a sudden evolution of gases. It is a mixture of saltpetre, charcoal, and sulphur. When heated, the saltpetre gives off oxygen and nitrogen; the oxygen combines with the charcoal, forming carbon dioxide and carbon monoxide; and the sulphur combines with the potassium, forming potassium sulphide. When a mixture of saltpetre and charcoal is burned, the reaction which takes place is this : 4KNO, + 50 = 3C0 2 + 2N a + 2K a CO s . By adding the necessary quantity of sulphur the carbon dioxide, which would otherwise remain in combination with the potassium as potassium carbonate, is given off, and potassium sulphide formed : 2KN0 3 + 30 + S = 3CO a + N a + K a S. For this reaction the constituents should be mixed in the proportions : Saltpetre 74.83 Charcoal 13.31 Sulphur. 11.86 100.00 This is approximately the composition of all powder. When gunpowder explodes, the gases formed occupy about 280 times the volume occupied by the powder itself. Potassium Nitrite, KNO a , is formed simply by heating the nitrate to a sufficiently high temperature. The reduc- tion is, however, much facilitated by adding to the nitrate some easily oxidized metal, as lead or iron. When the gases formed by the action of arsenic trioxide on nitric 464 COLLEGE CHEMISTRY. acid are passed into potassium hydroxide, both the nitrite and nitrate are formed, and they can be separated by crystallization. Potassium Chlorate, KOIO,. — The character of the reaction by which potassium chlorate is formed when chlorine acts upon a solution of potassium hydroxide has already been discussed (see p. 147). In the manufacture of the chlorate it is found advantageous to make calcium chlorate, and then to treat this with potassium chloride, when, at the proper concentration, potassium chlorate crystallizes out, on account of the fact that it is less solu- ble than the salts which are brought together. The process in brief consists in passing chlorine into a solution of calcium hydroxide in which an ercess of hydroxide is held in suspension. The first action consists in the formation of calcium hypochlorite. When the solution of this salt is boiled it is decomposed, yielding the chlorate and chloride: 3Ca(001), = Ca^Cl), + 20a€Bi. On now treating the solution with potassium chloride the following reaction takes place: Ca(O s Cl) 2 + 2KC1 = 2KC10 3 + Ca01 2 . Potassium chlorate crystallizes in lustrous crystals of the monoclinic system. Its taste is somewhat like that of saltpetre. It melts at a, comparatively low temperature (334°), and at 352° begins to decompose, with evolution of oxygen. At ordinary temperatures 100 parts of water dissolve 6 parts of the salt, and at the boiling temperature 60 parts. In consequence of the ease with which it gives up its oxygen, the chlorate is an excellent oxidizing agent, and it is constantly used in this capacity in the laboratory. Its oxidizing action is well illustrated by grinding a very little of it in a mortar with a little sulphur,* when an * Great care should be taken with all experiments with potassium chlorate. See description of experiments. POTASSIUM CYANIDE. 4^5 explosion takes place. With phosphorus the action is exceedingly violent. The chief uses of potassium chlorate are in the prepara- tion of oxygen, and in the manufacture of matches and fire- works. The tips of Swedish safety matches are made of potassium chlorate and antimony sulphide. The surface upon which they are rubbed to ignite them contains red phosphorus. The chlorate is frequently used in medicine, particularly as a gargle in cases of sore throat. Potassium Perchlorate, KC10 4 , is formed in the first stage of the decomposition of the chlorate by heat, as was explained under Oxygen (which see). It is prepared best by heating the chlorate in an open vessel until, after having been liquid, it begins to get solid again. Potassium Cyanide, KON. — Under Cyanogen it was stated that when nitrogen is passed over a highly-heated mixture of carbon and potassium carbonate, potassium cyanide is formed; and that carbon containing nitrogen compounds, as animal charcoal, when ignited with potas- sium carbonate, reduces the carbonate, forming potassium, in presence of which the carbon and nitrogen combine, forming the cyanide. The simplest way to make the cyanide is by heating potassium f errocyanide, K 4 Fe (CN) 6 , with potassium : K 4 Fe(C]Sr) 6 + 2K = 6KCN + Pe. Potassium cyanide is extremely easily soluble in water, and is deliquescent in moist air. When boiled with water it is decomposed, forming potassium formate and am- monia : KCN + 2H a O = KC0 2 H + NH,. It combines readily with oxygen when in the molten con- dition, as shown in its action upon lead oxide (see p. 380). In consequence of this power to combine with oxygen to form the cyanate, it is a valuable reducing agent, and is not infrequently used in the laboratory in this capacity. 466 COLLEGE CHEMISTRY. Potassium cyanide combines with the cyanides of the metallic elements, forming the so-called double cyanides of which potassium ferrocyanide, K 4 Fe(CN) 6 , is a good example. Potassium Sulphocyanate, KONS. — Just as potassium cyanide takes up oxygen to form the cyanate, it also takes up sulphur to form the sulphocyanate : KCN + S = KCNS. It is easily prepared by adding sulphur to molten potas- sium cyanide, or by heating a mixture of dehydrated potassium ferrocyanide, potassium carbonate, and sulphur. It crystallizes particularly well out of its solution in alcohol. It is deliquescent, and when dissolved in water it causes a very considerable lowering of the temperature. Thus, when 500 grams of the salt are mixed with 400 grams of water at the ordinary temperature, the tempera- ture sinks to about — 20°. Unlike the cyanate, it is not decomposed by water. Potassium Sulphate, KjSO^ — This salt occurs in com- bination with others in nature, particularly in the mineral kainite, which contains the constituents of potassium sul- phate, magnesium sulphate, and magnesium chloride, as indicated in the formula K 2 S0 4 .MgS0 4 .MgCl 2 + 6H 2 0. This occurs in Stassfurt and in Kalusz. Potassium sul- phate is used in medicine, and in the preparation of ordinary alum and of potassium carbonate. Primary, or Acid, Potassium Sulphate, KHS0 4 . — This salt is obtained as a secondary product in the preparation of nitric acid by the action of sulphuric acid upon salt- petre. It occurs in nature in the Grotto del Solfo, near Naples; and is made by treating the neutral salt with con- centrated sulphuric acid. "When heated above its melting- point it gives off water, and is transformed into the disulphate, thus: 2KHSO, = K 2 S 2 7 + H 2 0. When the disulphate is heated in contact with basic oxides POTASSIUM SULPHITES— CARBONATES. 467 it breaks down, forming sulphates. The decomposition is that represented in the equation K a S 2 0, = K 3 S0 4 +S0 3 . The nascent sulphur trioxide thus set free acts with great energy upon the oxides which are present. Hence acid potassium sulphate is a valuable agent for the purpose of decomposing some mineral substances which do not readily yield to the ordinary reagents. Its action consists in breaking down into the disulphate and water, the disul- phate then further breaking down into normal sulphate and sulphur trioxide. Sulphites. — "When sulphur dioxide is passed into a solu- tion of potassium carbonate until carbon dioxide ceases to escape, potassium sulphite, K 2 S0 3 , is formed. If the gas is passed to saturation the product is the primary or acid sulphite, KHS0 3 . Carbonates. — The normal salt, K 2 C0 3 , is the chief con- stituent of wood-ashes. When these are extracted with water the carbonate passes into solution and the salt thus obtained can be purified in a number of ways. The impure salt is known as potash. Formerly all the potassium car- bonate made was obtained from wood-ashes, but at present not more than half of the supply comes from this source. The other sources are the residues from the manufacture of beet-sugar, potassium sulphate and chloride, and wool- fat. The preparation of the carbonate from the sulphate and chloride is accomplished by the same method as that used in the preparation of sodium carbonate from the chloride. The methods used for this purpose will be treated of under the head of Sodium Carbonate (which see). The salt crystallizes from yery concentrated solu- tions in water. It is deliquescent, and dissolves in water with an evolution of heat, and the solution has a strong alkaline reaction. The explanation of the alkaline reaction of potassium carbonate in terms of the theory of electrolytic dissociation 468 COLLEGE CHEMISTRY. is very interesting. When the salt is dissolved in water it is dissociated into ions thns: K 2 C0 3 = 2K + 00,. With water, which is itself partly dissociated, the ions C0 3 combine to some extent to form the ions HC0 3 , while at the same time hydroxyl ions are left in the solution as shown in the equation : 00 3 + H 2 = HC0 3 + OH. But hydroxyl ions are believed to be the cause of basic reactions, so that the solution shows an alkaline -reaction. It will be observed that the cause of the phenomenon is what we call the weakness of the carbonic acid. A weak acid is only slightly dissociated in water solution. If the anions of the acid are formed in the presence of water they unite with hydrogen ions to form undissociated acid molecules, and this necessitates the formation of hydroxyl ions. These latter, however, do not combine with the cations K which are in the solution, as the product of this union would be potassium carbonate which is dissociated by the water. -Acid Potassium Carbonate, HK00 3 , is formed by pass- Zing, carbon dioxide over the normal salt, or into the con- centrated aqueous solution of the latter. It is much less easily soluble in water than the normal salt. The dry salt l gives off carbon dioxide and water easily when heated, and is converted into the normal salt : 2KHC0 3 = K 2 C0 3 + C0 2 + H s 0. The same decomposition takes place when the water solu- tion is heated, and even on evaporation at the ordinary temperature. Phosphates. — Three phosphates of potassium are known : (1) Tertiary, or normal potassium phosphate, K a P0 4 ; (2) RUBIDIUM— CAESIUM. 469 secondary, or di-potassium phosphate, K 2 HP0 4 ; and (3) primary, or mono-potassium phosphate, KH 2 P0 4 . There is nothing particularly characteristic about these salts, except the decompositions which the primary and secon- dary salts undergo when heated. These decompositions have already been referred to (see p. 450 and p. 451). Potassium Silicate, K 2 SiO s . — A compound of the definite composition represented by the formula here given has not been prepared. A solution of potassium silicate in water is prepared by dissolving sand or amorphous silicon dioxide in .potassium carbonate or hydroxide. It is prepared on the large scale by melting together quartz powder and purified potash. It. is known as water glass, for the reason that its solution dries in the air, forming a glass-like look- ing mass. To distinguish it from the water glass made with a sodium carbonate or hydroxide it is called potash tvater glass. Kubidiitm, Eb (At. Wt. 85.4). Cesium, Os (At. Wt. 133). Both these elements are widely distributed, but only in small quantities. They generally occur in company with potassium, which they resemble closely. They were dis- covered by means of the spectroscope by Bunsen and KirchhofE. The characteristic spectrum of rubidium con- sists of two dark red lines, and this is the origin of the name rubidium (from rubidus, dark red). Caesium was found in the Durkheim mineral water, and was recognized by two characteristic blue lines, and the name caesium was given to it on this account (from cwsius, sky-blue). Eubidium is found in different varieties of mica, known as lepidolite. The mineral pollux, which is essentially a silicate of caesium and aluminium, contains caesium as one of the chief constituents. It is a remarkable fact that the elements rubidium and caesium which are so similar to potasssium accompany it so generally in nature. Similar facts were noted in the 47° COLLEGE CHEMISTRY. group consisting of chlorine, bromine, and iodine, and that of sulphur, selenium, and tellurium. It will be remembered that chlorine is frequently accompanied by bromine and iodine; and sulphur by selenium and tel- lurium; but that chlorine and sulphur are present in much larger quantities than the elements which accompany them. Further, the relations between the atomic weights of the members of each group are approximately the same. Rubidium is prepared by the same method as that used in the preparation of potassium. It is silver-white with a yellowish tint. It can be con- verted into vapor which has a blue color. It takes fire in the air at the ordinary temperature. Its action upon water is the same as that of potassium, and its salts are very similar to those of potassium. Caesium has been prepared by the electrolysis of molten caesium cyanide. The salts of caesium are much like those of rubidium and potassium. y^ Sodium, Na (At. Wt. 23.05). Occurrence. — Sodium occurs very widely distributed and in large quantities in nature, principally as sodium chlo- ride. It is found in a number of silicates, and is a con- stituent of plants, especially of those which grow in the neighborhood of the seashore and in the sea. Just as the ashes of inland plants are rich in potassium carbonate, so the ashes of sea-plants and those which grow near the sea are rich in sodium carbonate. It is found everywhere in the soil, but generally in small quantities. Its presence in the soil is due to the decomposition of minerals con- taining it, such as soda feldspar, or albite. It occurs also as sodium nitrate or Chili saltpetre, and in large quantity in Greenland in the form of cryolite, which, as has been explained, is a so-called double fluoride of aluminium and sodium, of the formula ]Sra s AlP , or AlF,.3NaP. SODIUM. 471 r Preparation. — It is prepared by the action of an elec- ' trie current upon sodium hydroxide, or upon sodium chloride. Properties. — The properties of sodium are very similar to those of potassium. It is light, floating on water; it has a bright metallic lustre; and at the ordinary tempera- ture it is soft like wax. It decomposes water, but not as readily as potassium does. Its specific gravity is 0. 9735 ; its melting-point 95.6°. Its vapor is colorless when seen in thin layers, while thick layers appear purple. When melted and allowed to cool it takes the crystallized form. When exposed to the air it acts upon the moisture, and is converted into the hydroxide. Applications. — It is used for the purpose of isolating some elements whose oxides cannot easily be reduced, as, for example, aluminium, magnesium, and silicon, which are prepared by treating their chlorides with sodium. Silicon, however, as we have seen, is prepared better by treating potassium fluosilicate, K 2 SiF 6 , with sodium. The element is also used, in combination with mercury as sodium amalgam, a substance which affords a ready means of making nascent hydrogen. It also finds constant appli- cation in the laboratory for a variety of purposes. Sodium Hydride, Na 2 H, is formed in the same way ag the corresponding compound of potassium, and is in every way similar to it. Sodium Chloride, NaOl, — This is the substance which is generally known simply as salt, or common salt. It occurs very widely distributed, and in immense quantities, In the earth. The most important deposits are those at Wieliczka in Galicia, at Stassfurt and Eeichenhall in Germany, and at Cheshire in England. Besides these there are, however, many other deposits in the United States of America, in Africa, and in Asia. As it is easily soluble in water, many springs and streams, as well as lakes and the ocean, contain it. Sea-water contains 2.7 per cent. In some places sodium chloride is taken out of mines in solid form. Frequently, however, water is 472 COLLEGE CHEMISTRY. allowed to flpw into cavities in the earth, and to remain for some time in contact with the salt. The solution thus formed is afterward drawn or pumped out of the mine and evaporated by appropriate methods. It is generally allowed slowly to run down walls made of twigs, so that a large surface of the liquid is exposed to the air. The concen- trated solution thus obtained is then evaporated to crystal- lization by the aid of heat. In hot countries salt is obtained by the evaporation of sea-water, the heat of the sun being used for the purpose. Large shallow cavities are made in the earth, and into these the water flows at high-tide, or it is pumped up into them if they are too high. The process is continued for some months, and then the mother-liquor is drawn off, and the accumulated salt collected and subjected to proper methods of purification. The salt obtained by the above methods is not pure. It always contains sodium sulphate, together with magnesium and calcium chlorides. The chlorides of magnesium and calcium being deliquescent cause it to become moist in the air. Pure salt is not deliquescent. Sodium chloride crystallizes in colorless, transparent cubes. Some of that which occurs in nature has a blue color. When deposited from an evaporating solution it takes the form of small cubes arranged in groups of the shape of hollow pyramids, known as the hopper-shaped deposits. When deposited, the crystals enclose water, not as water of crystallization, and this is given off when the crystals are heated, the action being accompanied by a crackling sound. This is known as decrepitation. Sodium chloride melts at 815°, and is volatile at a red heat. In hot water it is but little more soluble than in cold. At 100° 100 parts of water dissolve 39 parts, and at ordinary temperatures 36 parts. Sodium chloride is the starting-point in the preparation of all sodium compounds, as well as of chlorine and hydrochloric acid. Salt is necessary to the life of man and many other animals. The r61e played by it in the SODIUM COMPOUNDS. 473 animal economy is not understood, but it is found generally distributed throughout the body in small quantity. The fluoride, bromide, and iodide of sodium are like the corresponding potassium salts and need not be described. Sodium Hydroxide, NaOH. — This compound resembles potassium hydroxide in all respects. Being cheaper it is used more extensively. It is prepared in the same way, by treating sodium carbonate in solution with calcium hydroxide, when insoluble calcium carbonate and soluble sodium hydroxide are formed : Na 2 CO, + Ca(OH) 2 = CaC0 3 + SNaOH. It is now prepared, for the most part, by the electrolysis of sodium chloride. The substance is commonly called caustic soda. It is extensively used for the purpose of making soap from fats. Oxides. — Sodium forms two oxides, the monoxide, Na 2 0, and the peroxide, Na 2 2 . In this respect a differ- ence is noticed between sodium and potassium; the latter forming the compounds K 2 and K 2 4 . Sodium Peroxide, Na 2 2 , has acquired importance in the arts as a bleaching-agent. It is prepared by heating sodium in a current of dry air at a temperature of 300°. When heated to a high temperature it gives off oxygen. Water decomposes it, forming sodium hydroxide, and setting oxygen free. Sodium Sulphantimonate, BTa,SbS 4 , also kn,own as Schlippe's salt, is a particularly beautiful example of the salts of sulpho-acids. It is made, as its composition indi- cates, by dissolving antimony pentasulphide in a solution of sodium sulphide : Sb 2 S 5 + 3Na 2 S = 2]STa 3 SbS,. r Sodium Nitrate, NaNO s . — This compound occurs in large quantity in southern Peru on the border of Chili, and is known as Chili saltpetre. The natural salt con- tains, besides the nitrate, sodium chloride, sulphate, and 474 COLLEGE CHEMISTRY. iodate. Sodium nitrate is similar to potassium nitrate, but it cannot be used in place of the more expensive potas- sium salt in the manufacture of the finer grades of gun- powder, as it becomes moist in the air, and does not decompose quickly enough. It is used extensively in the manufacture of nitric acid, and also for the purpose of preparing ordinary saltpetre. The iodine contained in Chili saltpetre is now extracted on the large scale, and this forms an important source of iodine. , Sodium Sulphate, Na 2 S0 4 . — This salt was first made by Glauber, as it is now made, by the action of sulphuric acid on sodium chloride. It is commonly called Glauber's salt. It occurs in a number of natural waters, as in that of Friedrichshall and Carlsbad. It occurs, further, in solid form in small quantities in some localities. It is made in very large quantities in connection with the manufacture of soda, the first reaction in this process being that of sodium chloride upon sulphuric acid. It is also formed in the manufacture of nitric acid by the action of sulphuric acid on Chili saltpetre. It crystallizes in large, colorless, monoclinic prisms, which contain ten molecules of water. These crystals are formed, however, only in case the temperature of the solution is below 20° at the time they are deposited. The salt is most easily soluble in water at 33°; above this point the solubility decreases. If deposited from its solution at temperatures above 33° it is anhydrous. Sodium sulphate easily forms supersaturated solutions which crystallize rapidly if disturbed, if a small crystal of the salt is thrown into them, and if cooled down to — 8°. This phenomenon is frequently presented by salts, but it is shown in a particularly striking way by this one. When exposed to the air the salt loses its water of crystallization and crumbles to a white powder. This is the process already described as efflorescence (see p. 84). Sodium sulphate is used as a purgative in medicine, and in the laboratory for the production of cold artificially. A good freezing mixture is made by bringing it together SODIUM CARBONATE. 475 with concentrated hydrochloric acid. Sodium chloride is formed, and the water of crystallization of the sulphate takes the liquid form. This change from the solid to the liquid form is accompanied by a marked absorption of heat. Ice can be made in this way without difficulty. The chief uses of the sulphate are in the manufacture of sodium carbonate and of glass, as will be explained farther on. --7 Sodium Thiosulphate, Na 2 S 2 3 + 5H 2 0.— This is the /salt which is commonly called hyposulphite of soda. It is made on the large scale by treating caustic soda with -sul- phur, and conducting sulphur dioxide into the solution. It is also made by boiling a solution of sodium sulphite and adding sulphur: Na 2 S0 3 + S = Na 2 S 2 3 . Its chief application is in photography, in which art it is used for the purpose of dissolving the excess of silver salt on the plate which has been exposed to the light, and on which a picture has been developed. The action consists in the f ormation- of salts in which both sodium and silver are contained. These are soluble in water. 7 Sodium Carbonate, ]STa 2 C0 3 . — This salt, commonly called soda, is one of the most important of manufactured chem- ical compounds. The mere mention of the fact that it is essential to the manufacture of glass and soap will serve to give some conception of its importance. It is found in the ashes of sea-plants, just as potassium carbonate is found in the ashes of inland plants. Formerly, it was made entirely from plant-ashes, but we are no longer dependent upon this source for our supply of the salt, as two methods have been devised for preparing it from sodium chloride, with which nature provides us in such abundance. As these methods are of great importance, and are, further, very interesting applications of chemical principles, they will be described below. Properties. — Anhydrous sodium carbonate is a powder which is formed by heating the crystallized salt. It melts 476 COLLEGE CHEMISTRY. to a clear liquid when heated to a sufficiently high tem- perature. It dissolves in water very readily with evolution of heat. The action is, however, not as marked as in the case of potassium carbonate. "When the salt is deposited from a water solution it has the composition Na 2 CO s -f 10H 2 O. This salt, it will be observed, contains the same number of molecules of water of crystallization as sodium sulphate. Like this, too, it effloresces when exposed to the air. When heated it melts in its water of crystalliza- tion, and the salt Na 2 C0 3 + H 2 0, or (HO) 2 C(ONa) 2 , separates. This, however, loses water when heated higher, and is converted into the anhydrous salt. The conduct of the carbonate towards water at different temperatures is suggestive of that of the sulphate. Its maximum solu- bility is at temperatures between 33° and 70°. Above the latter point the solubility decreases. The crystals of sodium carbonate containing ten molecules of water of crystallization belong to the monoclinic. system. Applications. — Sodium carbonate is used in immense quantities in the manufacture of glass, and in the prepara- tion of caustic soda, which is used in the manufacture of soap. / The Le Blanc Process for the Manufacture of Sodium / Carbonate. — In the manufacture of soda the problem to I be solved is the conversion of sodium chloride into sodium / carbonate. The first method devised for this purpose is / that of Le Blanc. During the French Bevolution the ' supply of potash was cut off from France. This led the government to offer a prize for a practical method for manufacturing soda from common salt. The method proposed by Le Blanc at that time, and which, until recently, has been used almost exclusively involves three reactions: (1) The sodium chloride is converted into sodium sul- phate by treating it with sulphuric acid : 2NaCl + H,SO t = Na 2 S0 4 + 2HC1. SODIUM CARBONATE: LE BLANC PROCESS. 477 (2) The sodium sulphate thus obtained is heated with charcoal, which reduces it to sodium sulphide: Na 2 S0 4 + 20 = Na s S + 2C0 2 . (3) The sodium sulphide is heated with calcium car- bonate, when sodium carbonate and calcium sulphide are formed: Na 2 S + CaC0 3 = Na 2 C0 3 + OaS. The conversion, of the sulphate into the carbonate is, therefore, expressed by the equation Na 2 S0 4 + 20 + CaCO, = Na/JO, + OaS + 2C0 3 . Calcium sulphide is insoluble in water containing lime, so that by treating the resulting mass with water the sodium carbonate is separated from the sulphide. In practice the sodium sulphate is mixed with coal and calcium carbonate, and the mixture heated in appropriately constructed furnaces. The coal reduces the sulphate to sulphide, which then reacts upon the calcium carbonate as above represented. The product of the action is known as crude soda or black ash. It contains, as its chief con- stituents, sodium carbonate and calcium sulphide, together with some calcium oxide, and a number of other substances in small quantities. In order to purify this product, it is broken to pieces, and treated with water; and the solution thus obtained evaporated, when the salt of the composition Na 2 C0 3 + 2H 2 is deposited. This is dipped out, and dried by heat, when it loses all its water. The product is the calcined purified soda of commerce. This always con- tains some sulphate and chloride togther with a small quantity of sulphite. When dissolved in water and allowed to crystallize, the salt is deposited in large crystals which contain water in the proportion represented by the formula ]STa 2 C0 3 -f- 10H 2 O. This is the so-called crystallized soda. 47 8 COLLEGE CHEMISTRY. As, in the manufacture of soda, by the Le Blanc process, the sulphur remains in combination as calcium sulphide, a process, known as the Chance process, has been devised for its recovery. This consists in passing carbon dioxide into the waste, thus liberating hydrogen sulphide; passing this into another portion of the waste, thus converting the calcium sulphide into the hydrosul- phide; and then treating this with carbon dioxide, when a gas rich in hydrogen sulphide is given off: C0 2 + CaH 2 S 2 + H 2 = CaCO s + 2H 2 8. By regulating the supply of air the gas is burned either to sulphur dioxide or to sulphur. / Ammonia Process for the Manufacture of Soda. — i Another process now in extensive use for the manufacture f of soda is the so-called ammonia process, or the Solvay process. This depends upon the fact that mono-sodium carbonate, HNaC0 3 , is comparatively difficultly soluble in water. If, therefore, mono-ammonium carbonate, or acid ammonium carbonate, HNH 4 C0 3 , is added to a solution of common salt, acid sodium carbonate, HNaC0 3 , crystal- lizes out, and ammonium chloride remains in the solution : NaCl + HNH 4 C0 3 = HNaC0 3 + NH 4 C1. "When the acid carbonate thus obtained is heated, it gives off carbon dioxide, and is converted into the normal salt thus: 2HNaC0 3 = tfa 2 C0 3 + C0 2 + H 2 0. i \ The carbon dioxide given off is passed into ammonia, and thus again obtained in the form of acid ammonium car- bonate : "o 1 NH 3 + H 2 + C0 2 = HNH 4 C0 3 . The ammonium chloride obtained in the first reaction is treated with lime or magnesia, MgO, and the ammonia set free. This ammonia is used again in the preparation of acid ammonium carbonate. The object of using mag- SODIUM PHOSPHATES. 479 nesia is to get magnesium chloride, which, when evaporated to dryness and heated, yields magnesia and hydrochloric acid: MgCl 2 + H 2 = MgO + 2HC1. A large proportion of the soda supply of the world is now furnished by the Solvay process. Mono-Sodium Carbonate, Primary Sodium Carbonate, HNaOO s . — This salt is commonly called " bi- carbonate of soda." It is easily prepared by passing carbon dioxide over the ordinary carbonate dissolved in its water of crys- tallization : Na^OO, + C0 2 + H 2 = SHNaCO When heated it gives up carbon dioxide and water, and is converted into the normal salt. As was stated in connec- tion with the ammonia-soda process, primary sodium carbonate is much more difficultly soluble in water than the normal salt. At ordinary temperatures 100 parts of water dissolve about 10 parts of the salt. It is used in medicine, and extensively in the prepara- tion of " soda water" and other effervescing drinks. Phosphates. — -There are three phosphates of sodium just as there are three phosphates of potassium. The point of chief interest presented by them is that the secondary salt, HNa 2 P0 4 , is the one most easily obtained, and is the substance commonly known as sodium phos- phate. When a solution of this salt is treated with an excess of sodium hydroxide, and the solution evaporated, normal or tertiary sodium phosphate crystallizes out. This has the composition N^PO, + 12H 2 0. The solution of the latter salt has an alkaline reaction, and when exposed to the air it absorbs carbon dioxide, and is converted into the secondary salt : 2Na 3 P0 4 + C0 2 + H 2 = 2H]Sra 2 P0 4 + Na 2 C0 8 . Secondary sodium phosphate, HNa 2 P0 4 + 12H 2 0, is easily made by adding sodium carbonate to a solution of 480 COLLEGE CHEMISTRY. phosphoric acid until an alkaline reaction is reached. It is also prepared on the large scale from bone-ash. It forms monoclinic prisms which effloresce in the air. Sodium Borate. — Normal boric acid, as we have seen, has the composition B(OH) s , and there are a number of borates derived from this acid. The salt which boric acid most readily forms with sodium hydroxide or sodium car- bonate, however, is that derived from tetraboric acid, H,B 4 0, , which is derived from normal boric acid by elim- ination of water. (See p. 34=2. ) This salt is borax, which in crystallized form has the composition represented by the formula Na 2 B 4 0, + 10H 2 O. Borax occurs in nature in several lakes in Asia and in Clear Lake, Nevada, in the United States. It is manu- factured by neutralizing, with sodium carbonate, the boric acid found in Tuscany. When heated, borax puffs up, and at red heat it melts, forming a transparent, colorless liquid. The dehydrated salt is known as anhydrous or calcined borax. In the molten condition, borax has the power to combine with metallic oxides, and, as many of the double borates thus formed are colored, the salt is used in blowpipe work for -the purpose of detecting certain metals. As it dissolves metallic oxides, it is used in the process of soldering, as it is necessary to have bright, un- tarnished metallic surfaces in order that the solder shall adhere firmly. The action of molten borax upon metallic oxides is similar to that which takes place when sodium hydroxide acts upon a solution of borax. Borates of the metals are formed together with sodium borate, or double borates in which part of the hydrogen is replaced by sodium and part by other metals. Borax is extensively used in the manufacture of porcelain and in glass-painting. It is an antiseptic, preventing the decomposition of some organic substances. -'Sodium Silicate, Na a Si0 3 .— Sodium silicate is formed by dissolving silicon dioxide in sodium hydroxide, and can be obtained in crystallized form. It is prepared on the large scale by melting together quartz sand and sodium LITHIUM— AMMONIUM SALTS. 481 carbonate in the proper proportions, and by melting together sodium sulphate, quartz sand, and charcoal powder. This substance is commonly known as water- glass. It is soluble in water, and, when its solution dries, it leaves a transparent coating on the surface on which it is placed. It is extensively used in the manufacture of artificial stone. Lithium, Li (At. Wt, 7.03). Lithium occurs in nature in relatively small quantity, chiefly in the minerals lepidolite, petalite, and spodumene, and in many mineral waters. It is also found in the ashes of a number of plants. It is prepared by the electrolysis of the chloride in the molten condition. The metal is silver-white, and is characterized by its low specific gravity. It acts vigorously upon water, but, if the water is at the ordinary temperature, the hydrogen given off does not take fire. In the air it conducts itself in much the same way that sodium does. The most characteristic salts of lithium are the phos- phate, carbonate, and chloride. Ammonium Salts. Attention has already been called to the marked simi- larity of the salts of potassium and sodium to those formed by the action of ammonia on the acids, and known as ammonium salts. The most important of these salts will be briefly treated of in this connection. A characteristic property of ammonium salts which distinguishes them from the salts of all the metals is their volatility. When sublimed, they all undergo decomposition, which is either partial or complete. The simplest kind of decomposition which they undergo is dissociation into ammonia and the acid. This is illus- trated in the case of ammonium chloride, which, when 482 COLLEGE CHEMISTRY. heated to a sufficiently high temperature, is dissociated into ammonia and hydrochloric acid. This is an example of true dissociation. The amount of decomposition is constant for any given temperature and pressure. • An ammonium salt of a polybasic acid containing some metal gives off ammonia and leaves an acid salt, which generally undergoes further decomposition. Thus, sodium- ammonium sulphate/NaNHjSO^, first gives off ammonia and forms mono-sodium sulphate: The acid salt thus formed then undergoes further change and the pyrosulphate is formed : 2S0 4 J g a = Na 2 S 2 0, + H 2 0. Another example of this kind of decomposition of am- monium salts is that afforded by sodium-ammonium phos- phate, HNaNH 1 P0 4 . When heated, this gives off ammonia and then water, the final product being sodium metaphos- phate : r ( ONa ( ONa PO i ONH, = PO i OH + NH S ; (OH (OH (ONa PO^OH = P0 2 ONa + H 2 0. ( OH Some ammonium salts undergo deeper-seated decom- positions, and do not give ammonia as one of the products. This is true especially of such salts as readily give off oxygen. In such cases the ammonia is oxidized, so that the hydrogen forms water. This is illustrated in the decomposition of ammonium nitrate and nitrite: NH 4 N0 3 = ]ST 2 + 2H 2 0; and NH 4 NO s = N a + 2H 3 0. AMMONIUM CHLORIDE— AMMONIUM SULPHIDE. 483 Further, all ammonium salts are decomposed with evolu- tion of ammonia when treated with basic hydroxides. This has been illustrated in the preparation of ammonia from ammonium chloride by treatment with calcium hydroxide : 2NH 4 01 + 0a(OH) a = Ca01 a + 2NH S + 2H a O. The ammonium salts are made by neutralizing acids with ammonia. sp,- Ammonium Chloride, NH 4 C1. — This salt is commonly if called sal ammoniac. At present its principal source is the so-called ammoniacal liquor of the gas-works. This liquid contains a considerable quantity of ammonium car- bonate, and, when it is treated with lime, ammonia is given off. This is passed into hydrochloric acid, and the solution of ammonium chloride thus formed evaporated to crystallization. The salt has a sharp, salty taste, and is easily soluble in water. "When heated, it is converted into vapor without melting and with very slight decomposition; and when the vapor comes in contact with a cold surface, it condenses in crystalline form. This process of vaporiz- ing and condensing a solid is called sublimation. Some of the ammonium chloride met with in the market has been sublimed. The salt is used in the preparation of ammonia, in medicine, and for other purposes. When it is dissolved in water, a considerable lowering of tempera- ture is caused. ~ Ammonium Sulphide, (NH 4 ) 2 S. — This compound is extensively used in chemical analysis for the purpose of precipitating those sulphides which are soluble in dilute hydrochloric acid (see p. 214 and p. 215). As will be remembered, in the usual method of analyzing a mixture of substances, the first step consists in adding hydrochloric acid to the solution. This precipitates silver, lead, and, under certain conditions, mercury. This precipitate having been filtered off, hydrogen sulphide is passed through the filtrate, when those metals are precipitated 484 COLLEGE CHEMISTRY. whose sulphides are insoluble in dilute hydrochloric acid. The precipitate is filtered off, and ammonium sulphide added to the filtrate, when the metals whose sulphides are soluble in dilute hydrochloric acid are thrown down. Among these are iron, cobalt, nickel, manganese, etc. Any other soluble sulphide might be used; but the advan- tage of ammonium sulphide is that it is volatile, and hence, by evaporating the solution and heating, it .can be gbt rid of after it has served its purpose. Another use to which it is put in analysis is for the purpose of dissolving the sulphides of tin, arsenic, and antimony, which are precipitated by hydrogen sulphide, and thus separating these from the other sulphides of the group. This solu- tion depends upon the power of the sulphides to form salts of sulpho-acids, as has been repeatedly explained. Ammonium sulphide is made by passing hydrogen sul- phide into an aqueous solution of ammonia. If the gas is passed until the solution, is saturated, the product is the hydrosulphide : NH 3 + H 3 S = NHJBS. If only half this quantity of the gas is passed, the product is the sulphide: 2NH 3 + H 2 S = (NH 4 ) 2 S. The simplest way to make it, however, is to divide a quantity of a solution of ammonia into two equal parts ; saturate one half, thus forming the hydrosulphide, and add the other half, when this reaction takes place: HNH 4 S + NH 3 = (NH 4 ) 2 S. The product is a colorless liquid of a disagreeable odor. It soon changes color, becoming yellow, and after a time a yellow deposit is formed in the vessel in which it is con- tained. This change of color is due to the action of the oxygen of the air. Some of the sulphide is decomposed into ammonia, water, and sulphur, thus: (NHJ.S + = 2NH, + H„0 + S. SODIUM-AMMONIUM PHOSPHATE. 485 The sulphur set free in this way combines with the nn- decomposed ammonium sulphide, forming the compounds (NH 4 ) 2 S 2 , (NHJ,S„ (NHJ,S 4 , and (NH 4 ) 2 S 5 . When as much sulphur has been set free as is required to form the pentasulphide, further decomposition by the oxygen of the air causes a deposit of sulphur. Therefore, in bottles containing ammonium sulphide which are allowed to stand for a long time a deposit of sulphur is always found. A solution containing the polysulphides is called yellow ammonium sulphide. It is this which is used for the pur- pose of dissolving the sulphides of arsenic, antimony, and tin in analytical operations. As stated above, a solution of ammonium hydrosulphide, HNH 4 S, is made by passing hydrogen sulphide into a solu- tion of ammonia until no more is taken up. ■? Ammonium Nitrate, NH^NO,, is obtained in crystals, which are easily soluble in water. It is of use chiefly in the preparation of nitrous oxide. When heated suddenly to a high temperature it is decomposed rapidly into nitrogen, water, and nitric oxide : 2NH 4 N0 3 = N 2 + 3NO + 4H 2 .^ \. ^ This decomposition may take place in the preparation of nitrous oxide if in the last stages of the operation the heat is raised too high, and explosions may be caused in this way. When dissolved in water a marked lowering of tem- perature takes place. Sodium-ammonium Phosphate, HNaNH 4 P0 4 . — This / salt is known as microcosmic salt, and is much used in the laboratory in blowpipe work. It is contained in guano and in decomposed urine. It is easily made by mixing solutions of di-sodium phosphate and ammonium chloride, and allowing to crystallize. In crystallized form it con- tains four molecules of water, HNaNH 4 P0 4 + 4H,0. The changes which the anhydrous salt undergoes when heated were described on page 482. When the crystallized salt i is heated, the water of crystallization is first given off. \ The value of the salt in blowpipe work depends upon the -^ 486 COLLEGE CHEMISTRY. fact that at high temperatures the metaphosphate com- bines with metallic oxides, forming mixed phosphates, the reactions being like those which metaphosphoric acid undergoes with water: 2HPO, + H 2 = H 4 P 2 7 ; , HPO, + H 2 = H 3 P0 4 ; 2NaP0 3 + M 2 = Na 2 M 2 P 2 0,; NaPO, + M 2 = NaM a P0 4 . Many of these double phosphates and pyrophosphates are colored, and furnish a means of detecting some of the metals. Reactions of the Members of the Sodium Group which are of Value in Chemical Analysis. — The chief difficulty experienced in chemical analysis is in distinguishing between similar elements. Sodium and potassium, for example, conduct themselves so much alike in so many respects that we might subject them to the influence of a number of reagents without being able to tell which one we are working with. For purposes of analysis, therefore, it is necessary to take advantage of differences between the elements, and the more striking the differences the better. Those reactions which give rise to the formation of insolu- ble compounds or precipitates are most frequently used in analysis. Very few salts of the members of the sodium group are insoluble, and the difficulty of distinguishing between these elements is increased by this fact. In ordinary analyses the elements of this group which are of most importance are potassium and sodium, the other elements of the group being but rarely met with. Am- monium compounds are easily distinguished from those of potassium and sodium by the fact that, when treated with caustic soda or potash, they give off ammonia, which is recognized by its characteristic odor. The chief reactions which are of value in distinguishing between potassium and sodium are the following : Chlorplatinic Acid, H 2 Pt01 4 , forms difficultly soluble salts with potassium and ammonium chlorides. These FLAME REACTIONS AND THE SPECTROSCOPE. 487 are the chlorplatinates, K 2 PtCl 6 and (NH 4 ) 2 PtCl 6 . The corresponding salt of sodium is easily soluble. Perchloric Acid, HC10 4 , forms difficultly soluble potas- sium per chlorate, KC10 4 , when added to solutions of potassium salts. Fluosilicic Acid, H 2 SiF 6 , forms difficultly soluble salts ■with potassium and sodium, K 2 SiF 6 and Na 2 SiF 6 , but not with ammonium. Tartaric Acid, H 2 (C 4 H 4 6 ), forms a difficultly soluble potassium salt of the formula KH(C 4 H 4 6 ). The corre- sponding salt of sodium is easily soluble. The formation of mono-potassium tartrate takes place as represented in the equation: KOI + H 2 (C 4 H 4 6 ) = KH(C 4 H 4 6 ) + HOL Normal or neutral potassium tartrate is soluble in water, so that, if the difficultly soluble acid tartrate is filtered off, and potassium carbonate added to it, it dissolves in consequence of the formation of the neutral salt, which takes place as represented in the equation 2KH(C 4 H 4 6 ) + K 3 C0 3 = 2K 2 (C 4 H 4 6 ) + CO a + H 2 0. If, to the solution of the neutral salt, hydrochloric acid is added, the acid salt is again formed and precipitated : K 3 (C 4 H 4 6 ) + HC1 = KH(C 4 H 4 6 ) + KOI. Di-sodium Pyro-antimonate, Na 2 H 2 Sb 2 T , is insoluble in cold water, and is formed when a solution of the corre- sponding potassium salt is added to a solution of a sodium salt. Flame Reactions and the Spectroscope When a clean piece of platinum wire is held for some time in the flame of the Bunsen burner, it then imparts no color to the flame. If now a small piece of sodium carbonate or any other salt of sodium is put on it, the flame is colored in- 488 COLLEGE CHEMISTRY. tensely yellow. All sodium compounds have this power, and hence the chemist makes use of this fact for the pur- pose of detecting the presence of sodium. Similarly, potassium compounds color the flame violet; lithium com- pounds color the flame red; rubidium and caesium produce colors similar to that of the potassium flame. While it is an easy matter to recognize potassium alone, or any one of the other metals alone, it is difficult to do so when they are together in the same compound. For example, when sodium and potassium are together, the intense yellow caused by the sodium completely masks the more delicate violet caused by the potassium, so that the latter cannot be seen by the unaided eye. In this particular case the difficulty can be got over by letting the light from the flame pass through a blue glass, or through a thin vessel of glass containing a solution of indigo. The yellow light is thus cut off, while the violet light passes through and can be recognized. A more general method for detecting the constituents of light is by means of a prism of glass. Lights of different colors, which are produced by ether waves of different lengths, are turned out of their course to different extents when passed through a prism, as is seen when white sunlight is passed through a prism. A narrow beam of white light passing in emerges as a band of various colors, called its spectrum. We thus see that white light is made up of lights of different colors; or, to speak in the language of physics, that motion of the light- ether which produces upon the eye the sensation of white light is made up of a number of motions, each of which alone produces upon the eye the sensation of a color. Similarly, we can determine what any light is composed of. Every light has its characteristic spectrum. The light given off from any solid heated to a white heat gives a continuous spectrum, like that of the sunlight. An in- candescent gaseous substance, on the other hand, gives a spectrum made up of separate bands of color, or a line spectrum. The light produced by burning sodium, or by introducing a sodium compound in a colorless flame, gives EXPERIMENTS WITH POTASSIUM SALTS. 4«9 a spectrum consisting of a narrow yellow band. The spectrum of the potassium flame consists essentially of two bands, one red and one violet. Further, these bands always occupy definite positions relatively to one another, so that, in looking through a prism at the light caused by potassium and sodium, the yellow band of sodium is seen in its position, and the two potassium bands in their proper positions. There is therefore no difficulty in detecting these elements when present in the same sub- stance or in the presence of other elements which give characteristic spectra. The instrument used for the purpose of observing the spectra of different lights is called the spectroscope.* It consists essentially of a prism and two telescopes. Through one of the telescopes the light to be examined is allowed to pass so as to strike the prism properly. The light emerges from the other side of the prism, and is observed through the other telescope, which is provided with lenses for the purpose of magnifying the spectrum. By means of a third telescope, an image of a scale is thrown upon the face of the prism from which the spectrum emerges, and is reflected thence into the observing-tube, together with the spectrum, so that the position of the bands can be accurately determined. By means of the spectroscope, it is possible to detect the minutest quantities of some elements, and, since it was devised, several new elements have been discovered through its aid; as, for example, caesium, rubidium, thallium, indium, gallium, and others. EXPERIMENTS. Potassium Salts. •\ Experiment 209. — In preparing potassium iodide from iodine and potassium hydroxide, proceed as follows : To 30 grams iodine *For an account of the spectroscope and its uses, the student should consult some work on physics. The principles involved in its construction and application are physical principles, and cannot properly be taken up in detail in a text-book of chemistry. 49° COLLEGE CHEMISTRY. use 15 grams hydroxide. Dissolve the latter in 100 cc. water. Add half this solution to the iodine in a porcelain evaporating- dish. Now slowly add the rest of the liquid until the color dis- appears. Concentrate the liquid to a syrupy consistence, add 1 gram finely-powdered charcoal, mix, and evaporate to dryness. The residue is then heated to redness in an iron vessel. After cooling extract with water. Experiment 210. — Potassium iodide can also be prepared by the following method: Bring together in a capsule 200 grams water, 10 grams iron filings, and 40 grams iodine ; mix, and heat gently. When the solution has become green, decant, filter, and wash. Now heat the liquid nearly to boiling, and gradually add a solution of 35 grams potassium carbonate in 100 grams water. Filter, wash, and evaporate. Experiment 211. — Dissolve 50 grams potassium carbonate in 500 to 600 cc. water. Heat to boiling in an iron or a silver ves- sel, and gradually add the slaked lime obtained from 25 to 30 grams of good quicklime. During the operation the mass should be stirred with an iron spatula. After the solution is cool, draw it off by means of a siphon into a bottle. This may be used in experiments in which caustic potash is required. Experiment 212. — Mix together 15 grams potassium nitrate and 2. 5 grams powdered charcoal. Set fire to the mass. Experiment 213. — Treat a quantity of wood-ashes with water. Filter, and examine by means of red litmus paper. Evaporate to dryness. What evidence have you that the residue contains potassium carbonate ? Sodium Salts. --- Experiment 214. — Make a supersaturated solution of sodium sulphate by heating an excess of the salt with water at 33°. Filter the solution into small flasks and cork them. On remov- ing the corks and agitating the vessels, the salt will suddenly crystallize out. Experiment 215. — Make a saturated solution of common salt in ordinary ammonia- water (about 50 cc). Pass carbon dioxide into this solution until no more is absorbed, the delivery-tube being arranged as shown in Experiment 75. Filter off the pre- cipitate, and dry it by spreading it upon layers of filter-paper. Heat some of the salt when dry, and determine whether the gas given off is carbon dioxide or not. When gas is no longer given EXPERIMENTS WITH SODIUM SALTS. 49* off by heating, let the tube cool and examine the residue. [Is it a carbonate ?] Experiment 216. — Make ammonium sulphide thus : Divide a given quantity of a solution of ammonia into two equal parts. Saturate one half by passing hydrogen sulphide through it, and then add the other half, CHAPTEE XXVI. ELEMENTS OF FAMILY II, GROUP A: GLUCINUM.— MAGNESIUM— CALCIUM.— STRONTIUM. —BARIUM [ERBIUM]. General. — The elements of this group fall into two sub- groups. Calcium, strontium, and barium are strikingly alike. They also have some points in common with the members of the potassium family, and at the same time are related in some degree to the metals of Family III, Group A, which are known as the earth metals. There- fore, calcium, barium, and strontium are generally called the metals of the alkaline earths. Glucinum and magnesium resemble the metals of the alkaline earths in some ways, but they also resemble the members of Group B, of the same family, which includes zinc and cadmium. On com- paring the group with the elements presented in the last chapter, some analogous facts are noticed. Arranging the irve elements of the potassium group in the order of their atomic weights, and the elements of Family II, Group A, in the same way, we have this table : Li Na K Eb Cs 7.03 23.05 39.15 85.4 133 Gl Mg Ca Sr Ba 9.1 24.36 40 87.6 137.4 As regards the analogies between the elements in each group, the last three members of each group resemble one another more closely than they resemble the first two members of the group, while the first two members in each group also resemble each other closely. The natural 492 ELEMENTS OF FAMILY II, GROUP A. 493 grouping according to the properties is into the sub- groups : a b y.,-,. Potassium, and Eubidium, Sodium, Caesium. ^t . Calcium, Glucinum, ' , T . and Strontium, Magnesium, „ . ' 8 Barium. The relations between the atomic weights of the elements of Family II, Group A, are similar to those of the elements of Family I, Group A. That of magnesium, 24.36, is nearly half the sum of those of glucinum, 9.1, and cal- cium, 40. We have ?^ + i-° = 24.55. % So, also, that of strontium, 87.6, is approximately half the sum of those of calcium, 40, and barium, 137.4. 40 + 9 137 -* = 88.7. In the calcium group the specific gravities increase in the order of the atomic weights : At. Wt. Sp. Gr. Calcium 40 1.57 Strontium 87.6 2.5 Barium 137.4 3.75 All the elements of the group are bivalent. The general formulas of the principal compounds are as follows : MCI,, M(OH)„ M(N0 3 ) 2 , MS0 4 , M 3 (P0 4 ) 2 , MSiO s , etc. The chlorides, hydroxides, and nitrates are soluble in water. The sulphates decrease in solubility as the atomic weights increase. Glucinum sulphate, G1S0 4 , is soluble 494 COLLEGE CHEMISTRY. in its own weight of water; magnesium sulphate, MgS0 4 , is soluble in about three times its weight of water; calcium sulphate, CaS0 4 , dissolves in 400 parts; strontium sul- phate, SrS0 4 , in about 8000 parts; and barium sulphate, BaS0 4 , in about 400,000 parts of water. Barium sul- phate, as will be seen, is practically insoluble in water. The normal carbonates of all except glucinum are insoluble in water. The solubility of the hydroxides increases as the atomic weight increases. Glucinum hydroxide is in- soluble; magnesium hydroxide is but slightly soluble. One hundred parts of water at the ordinary temperature dissolve 0.1368 part of calcium hydroxide, 2 parts of strontium hydroxide, and 3.5 parts of barium hydroxide. The solubility of strontium and barium hydroxides is, however, much increased at higher temperatures. CALCIUM SUB-GROUP. This sub-group, as bas been stated, consists of the three similar elements, calcium, strontium, and barium. Of these calcium occurs most abundantly in nature. Barium and strontium frequently accompany each other, and both are found in some localities in company with calcium. They are much less abundant in nature than calcium. Calcium, Ca (At. Wt. 40). " ^ Occurrence. — Calcium is found in nature in enormous quantities, chiefly in the form of the carbonate, CaC0 3 , as limestone, marble, and chalk. It also occurs in the form of the sulphate, CaS0 4 , as gypsum; of the phos- phate, Ca 3 (P0 i ) 2 , as phosphorite and apatite; of the fluoride, CaF a , as fluor-spar. It is found in solution in most natural waters either as the carbonate or sulphate; and in the organs of plants and animals. Bones contain a large proportion of calcium phosphate; egg-shells and coral contain calcium carbonate. CALCIUM. 495 y? Preparation. — Calcium is made : (1) by heating together ' anhydrous calcium iodide and metallic sodium; (2) by heating calcium oxide and carbon together in an electric furnace in a current of hydrogen. In the latter case a mixture of calcium and the hydride is obtained, but the metal can be separated from this without difficulty. / Properties. — Calcium is a silver- white metal. It is soft /and can be cut with a knife like sodium. It is one of the most active elements known, and combines with all the other elements except argon. It combines with hydrogen to form the hydride CaH a , and burns brilliantly in the air, forming compounds with both oxygen and nitrogen. It decomposes water easily and becomes incandescent if in the form of powder. Calcium Chloride, Ca01 2 . — This salt is found in nature in combination with other chlorides, particularly in the mineral tachydrite, which occurs in the salt deposits at Stassfurt, and has the composition represented by the formula CaCl 2 . MgCl a + 12H a O. It is also found in solu- tion in sea-water. It is obtained as a by-product in the preparation of ammonia from ammonium chloride and lime; in the preparation of potassium chlorate from cal- cium chlorate and potassium chloride (see p. 464); and in the ammonia-soda process. It is made by dissolving calcium carbonate in hydrochloric acid, as in the prepara- tion of carbon dioxide. Prom very concentrated solutions it crystallizes with six molecules of water, CaCl 2 -f 6H 2 0. When these crystals are exposed to the air they soon deliquesce. When a solution of calcium chloride is evap- orated, and care is taken to keep the temperature below 200°, it solidifies, forming a porous mass which has the composition represented by the formula CaCl 2 -f- 2H s O. This is much used in laboratories as a drying agent, as it absorbs water with great ease. If this salt is heated above 200° it loses all its water, and the dehydrated chloride melts, forming fused calcium chloride. This is also much used on account of its drying power. Cases are passed through tubes filled with granulated calcium chloride for 49 6 COLLEGE CHEMISTRY. the purpose of drying them, and the salt is also placed in vessels in which it is necessary that the air should be dry, as in balance-cases, desiccators, etc. Calcium chloride forms crystallized compounds with ammonia and with alcohol, as well as with water. It is obvious from this that calcium chloride cannot be used for the purpose of drying ammonia gas. When the compounds with ammonia and with alcohol are heated they break down, yielding ammonia and alcohol respectively, as the compound with water gives up the latter. j» Calcium Fluoride, CaF 3 . — This compound occurs in large quantities in nature as the mineral fluor-spar. It occurs beautifully crystallized in cubes, and is insoluble in water. It is the source of fluorine compounds in general, and is used in metallurgical operations for the reason that it melts readily and does not act upon other substances easily. It therefore simply serves as a liquid medium in which reactions take place at high temperatures. A sub- stance that acts in this way and is used for this purpose is called a flux. The name fluor-spar has its origin in , this use of the substance. ^- Calcium Oxide, CaO. — This important compound is commonly called lime, or, to distinguish it from the hydroxide or slaked lime, it is called quick-lime. It is made in large quantity by heating calcium carbonate in appropriately constructed furnaces, known as lime-kilns. Pure lime is made by decomposing some pure form of calcium carbonate, as marble or calc-spar. Lime is a white, amorphous, infusible substance. When heated in the flame of the compound blowpipe it gives an intense light, as any other infusible substance would do under the same circumstances. When exposed to the air it attracts moisture and carbon dioxide, and is converted into the carbonate. It must therefore be protected from the air. Lime which has been converted into the car- bonate by exposure to the air is said to be air-slaked. -^ Calcium Hydroxide, Ca(OH) 2 . — When calcium oxide or quick-lime is treated with water it becomes hot and CALCIUM HYDROXIDE. 497 crumbles to a fine powder. The substance formed in this operation is somewhat soluble in water, the solution being known as lime-water. The chemical change that takes place when lime is treated with water has been explained. It consists in the formation of a compound of the formula Ca(OH) 2 , known as slaked lime; and the operation is known as slaking. The action is of the same kind as that with which we have so frequently had to deal in the trans- formation of oxides into the corresponding hydroxides. Thus when potassium oxide is treated with water it is changed to the hydroxide, with a marked evolution of heat, the reaction being represented in this way: K 2 + H 2 = |°g. So, too, when sulphur trioxide is brought in contact with water it appears to form the hydroxide, normal sulphuric acid: fOH S \ = + 3H 2 = S 4 }-=0 OH OH OH OH OH The action in the case of calcium oxide is represented in a similar way : Ca = + H 2 = Ca<^- The hydroxide is a fine white powder. At red heat it loses water and is reconverted into the oxide : Ca<°g = Ca = + H 2 0. When lime-water is exposed to the air it becomes covered with a crust of calcium carbonate, and finally all the cal- cium is precipitated as calcium carbonate. A solution of calcium hydroxide affords a convenient means of detecting 498 COLLEGE CHEMISTRY. the presence of carbon dioxide, as has been shown in deal- ing with this gas. The solutiqn has an alkaline reaction, and acts in many respects like solutions of the hydroxides of potassium and sodium. Applications. — Lime is extensively used in the arts, generally in the form of the hydroxide. As we have seen, it is used in the preparation of ammonia and the caustic alkalies, potassium and sodium hydroxides; and of bleach- ing-powder and potassium chlorate. It is further used in large quantity for the purpose of removing the hair from hides in the process of tanning; in decomposing fats for the purpose of making stearin for candles; for purifying gas; and especially in the preparation of mortar. >», Bleaching-powder. — The preparation of bleaching- powder was referred to under Chlorine (which see). The main reaction involved is that represented in the equation 2Ca(OH) 2 + 4C1 = Ca(C10) 2 + CaCl 2 + 2H 2 0. Bleaching-powder. The compound is commonly called "chloride of lime." Assuming that the reaction takes place in the same way a3 that of chlorine on caustic potash, the product is a mix- ture of calcium hypochlorite, Ca(C10) 2 , and calcium chloride, for it is held that the reaction with potassium hydroxide takes place as represented in this equation : 2KOH + 2C1 = KCIO + KC1 + H 2 0. An objection to the view that calcium chloride is present as such in bleaching-powder is found in the fact that the substance is not deliquescent, as it should be if calcium chloride were present. This has led to the suggestion that bleaching-powder in the dry form is not a mixture of two compounds, as represented above, but that it is rather one ( CI compound of the formula Ca ■] ^p, or CaOCl 2 . A com- pound of this formula would plainly have the same com- position as a mixture of calcium hypochlorite and calcium BLEACHING-POWDER. 499 chloride in the proportion of their molecular weights. For we have Ca(C10) 2 + CaCl 2 = 2CaOCl 2 . The point is a difficult one to decide, but at present the evidence appears to be rather in favor of the view that bleaching-powder in the dry form is a single compound of the constitution represented by the last formula given. "When treated with water, however, it conducts itself like a mixture of the hypochlorite and chloride. This is entirely in accordance with the theory of electrolytic dis- sociation, as the ions would be the same in both cases : CaCl 2 = Ca + 2C1; Ca(OCl) 2 = Ca + 20C1. 2Ca Ca. The salt of the formula CaS0 4 also occurs in nature, and is called anhydrite. A granular form of gypsum is called alabaster. Calcium sulphate is difficultly soluble in hot and cold water, but its solubility is markedly increased by the pres- ence of certain other salts; as, for example, sodium chloride. It is comparatively easily soluble in hydrochloric acid and in nitric acid. When heated to 100°, or a little above, it loses nearly all its water and forms a powder known as plaster of Paris, which has the power of taking up water and forming a solid substance. This process of solidification is known as "setting." Plaster of Paris is very largely used in making casts, on account of its power to harden after having been made into a paste with water. The hardening is a chemical process, and is caused by the combination of water with the salt to form the crystallized variety : CaS0 4 + 2H 2 = (HO) 4 SCa. "When heated to 200°, and above, all the water is given off from gypsum, and the product now combines with water only very slowly, and is of no value for making casts. In general, the higher the temperature to which the gypsum is heated, the greater the difficulty with which the product combines with water. Many natural waters contain gypsum in solution. Such waters act in some respects like those which contain cal- cium carbonate. With soap, for example, they form insoluble compounds. They are called hard waters. This kind of hardness is not removed % boiling, and it is S04 COLLEGE CHEMISTRY. therefore called permanent hardness. Magnesium sulphate acts in the same way, producing permanent hardness. When calcium sulphate is treated with a solution of a soluble carbonate, it is decomposed, forming the carbonate as represented in the equation CaS0 4 + m 2 G0 3 = Na 2 S0 4 + CaCO s . Thic change is effected by allowing the two to stand in contact at the ordinary temperature. Applications. — Besides being used for making casts, calcined gypsum is used in surgery for making plaster-of- Paris bandages, and as a fertilizer. Its action as a fertili- zer is believed by some to be due to the fact that it has the power to hold ammonia and ammonium carbonate in combination, and thus to make them available for the plants. It has recently been shown that it in some way facilitates the process of nitrification, and perhaps it is in v consequence of this that it facilitates plant-growth. V Calcium Phosphates. — There are three phosphates of | calcium: (1) The normal phosphate^ Ca 3 (P0 4 ) 2 ; (2) the I secondary phosphate, 43aHP0 4 ; and (3) the primary phos- \ jo7ia^CaH 4 (POJ 2 . ^ (1) Normal calcium phosphate, Ca s (P0 4 ) 2 , is derived from phosphoric acid by the substitution of calcium for all the hydrogen. It is found in nature in large quantity as phosphorite, and in combination with calcium fluoride or chloride as apatite. It is, further, the chief inorganic constituent of bones, forming 85 per cent of bone-ash, and is contained in the excrement of animals, as in guano, etc. It is found everywhere in the soil, and is taken up by the plants for whose development it is essential. That it is also essential to the life of animals is obvious from the fact that the bones consist so largely of it. The phosphate needed for the building up of bones is taken into the system with the food. From these statements, it is clear that calcium phosphate is of fundamental importance, and that a fertile soil must either contain this salt or something from which it can be formed. Now, when a crop is raised CALCIUM PHOSPHATES. JoJ on a given area, a certain amount of the phosphate con- tained in it is withdrawn. If the plants were allowed to decay where they grow, the phosphate would be returned and the soil would continue fertile; but in cultivated lands this is not the case. The crops are removed, and with them the calcium phosphate contained in them, and the soil therefore becomes exhausted. If the substances removed are used as food, some of the phosphate is found in the excrement of the animals; and, if this excrement is put on the soil, it is again rendered fertile. There are, however, other sources of calcium phosphate, and some of these are utilized extensively in the preparation of artificial fertilizers. The natural form of the phosphate, as that in bone-ash, in phosphorite, and in guano, is mainly the normal or neutral phosphate. This is insoluble in water, and is therefore taken up by the plants with difficulty. To make it quickly available, it must be converted into a soluble phosphate. This is done by treating it with sul- phuric acid in order to efEect the reaction represented in this equation : Ca 3 (P0 4 ) a + 2H 2 S0 4 = CaH 4 (P0 4 ) 3 + 2CaS0 4 . The primary phosphate thus formed is soluble in water, and is of great value as a fertilizer. The mixture of the soluble phosphate and of calcium sulphate is known as "superphosphate of lime." The sulphate, as we have seen, is also of value as a fertilizer. The value of super- phosphates depends chiefly upon the amount of soluble phosphate contained in them; and in dealing with them it is customary to state how much " soluble" and how much "insoluble phosphoric acid" they contain. When a superphosphate is allowed to stand for a time, some of the soluble primary phosphate is converted into insoluble phosphates by contact with basic hydroxides and water. This is known as the process of " reversion," and that part of the phosphoric acid which is contained in the insoluble phosphate is spoken of as "reverted phosphoric acid." Normal calcium phosphate, as has been stated, is in- 5°6 „ COLLEGE CHEMISTRY. soluble in water, and is formed when a soluble normal phosphate is added to a solution of a calcium salt. It is also formed when di-sodium phosphate and ammonia are added to a solution of a calcium salt, thus : 2HNa 2 P0 4 + 3CaCl a + 2NH 3 = Ca 3 (P0 4 ) a + 4NaCl + 2NH,C1. Normal or tertiary calcium phosphate is soluble in hydrochloric acid and in nitric acid, in consequence of the formation of calcium chloride, or nitrate, and the primary phosphate. If ammonia is added to this solution, the tertiary phosphate is again precipitated, as represented below : Ca 3 (P0 4 ) 2 + 4HC1 = 2CaCl 2 + H 4 Ca(P0 4 ) 2 ; 2CaCl 2 + H 4 Ca(P0 4 ) 2 + 4NH 3 = Ca 3 (P0 4 ) 2 + 4NH 4 C1. (2) Secondary calcium phosphate, CaHP0 4 , is formed, as above described, when a solution of a calcium salt is treated with secondary sodium phosphate. (3) Primary calcium phosphate, H 4 Ca(P0 4 ) s , is com- monly called the acid phosphate of calcium. It is formed when ordinary insoluble calcium phosphate is treated with concentrated sulphuric acid, and is contained in the so-called superphosphates. It is also formed by treating the neutral phosphate with phosphoric acid and with hydrochloric acid. When treated with but little water, it is converted into the secondary salt and free acid : H 4 Ca(P0 4 ) 2 = HCaP0 4 + H s P0 4 . Calcium Silicate, CaSi0 3 , occurs in nature as the mineral wollastonite, and, in combination with other sili- cates, in a large number of minerals, as garnet, mica, the zeolites, etc. It is formed when a solution of sodium sili- cate is added to a solution of calcium chloride, and when a mixture of calcium carbonate and quartz is heated to a high temperature. Glass. — Ordinary glass is a silicate of calcium and sodium made by melting together sand (silicon dioxide, SiO,) with lime and sodium carbonate or soda. Instead GLASS. 507 of calcium carbonate, lead oxide may be used ; and instead of sodium carbonate, potassium carbonate. The proper- ties of the glass are dependent upon the materials used iu its manufacture. Ordinary window glass is a sodium-calcium glass. The purer the calcium carbonate and silica, the better the quality of the glass. This glass is comparatively easily acted upon by chemical substances, and is therefore not adapted to the preparation of vessels which are to be used to hold acids and other chemically active substances. It answers, however, very well for windows. The difference between ordinary window glass and plate glass is essentially that the former is blown and then cut into pieces, while the latter, when in the molten condition, is run into flat moulds and there allowed to solidify. Bohemian glass is made with potassium carbonate. If pure carbonate is used, as well as pure calcium carbonate and silica, a very beautiful glass is the result. It is char- acterized by great hardness, by its difficult fusibility, and by its resistance to the action of chemical substances. It is -particularly well adapted to the manufacture of vessels and tubes for use in chemical laboratories. Flint-glass is made by melting together lead oxide, potassium carbonate, and silicon dioxide. It is character- ized by its power to refract light, its high specific gravity, its low melting-point, and the ease with which it is acted upon by reagents. Owing to its high refractive power, it is largely used in the manufacture of lenses for optical instruments. Strass is a variety of lead-glass that is particularly rich in lead. Its refracting power is so great that it is used in the manufacture of artificial gems. Colors are given to glass by putting in the fused mass small quantities of various substances. Thus, a cobalt compound makes glass blue; copper or chromium makes it green; one of the oxides of copper makes it red; uranium gives it a yellow color; etc. The most common variety of glass is that used in the manufacture of ordinary bottles. 5<>8 COLLEGE CHEMISTRY. It is generally green to black, and sometimes brown. In its manufacture, impure materials are used, chiefly ordi- nary sand, limestone, sodium sulphate, common salt, clay, etc. "When glass is suddenly cooled, it is very brittle and breaks into small pieces when scratched or slightly broken in any way. This is shown by the so-called Prince Eupert's drops, which are made by dropping glass, in the molten condition, into water. When the end of such a drop is broken off, the entire mass is completely shattered into minute pieces. It is clear from this that, in the manufacture of glass objects, care must be taken not to cool them suddenly. In fact they are cooled very slowly, the process being known as annealing. For this purpose they are placed in furnaces the temperature of which is but little below that of fusion, and they are kept there for some time, the heat being gradually lowered. If red-hot glass is introduced into heated oil or paraffin, and allowed to cool very slowly, it is found to be extremely hard and elastic. The glass of De la Bastie is made in this way. Vessels made of it can be thrown about upon hard objects without breaking, but sometimes a slight scratch will cause the glass to fly in pieces as the Eupert's drops do. Mortar. — Mortar is made of slaked lime and sand. When this mixture is exposed to the air, calcium carbonate is slowly formed and the mass becomes extremely hard. - The water contained in the mortar soon passes off, but nevertheless freshly plastered rooms remain moist for a considerable time. This is due to the fact that a reac- tion takes place between the carbon dioxide and calcium hydroxide by which calcium carbonate and water are formed, Ca(OH) a + C0 2 = CaCO s + H a O, and it is the water thus liberated that keeps the air moist. The complete conversion of the lime into car- bonate requires a long time, because the carbonate which MORTAR-CALCIUM SULPHIDE. 509 is formed on the surface protects, to some extent, the lime in the interior. It is generally regarded as unhealthy to live in rooms with freshly plastered walls, because the air is constantly kept moist in consequence of the reaction above mentioned. It is, however, difficult to see why the presence of a little extra moisture in the air should be unhealthy; and, if there is any danger from freshly plastered walls, it seems probable that the cause must be sought for elsewhere. It is possible that the constant presence of moisture in the pores of the wall interferes with the important process of diffusion, and that therefore when the room is closed this natural method of ventilation cannot come into play. When limestones which contain magnesium carbonate and aluminium silicate in considerable quantities are heated for the preparation of lime, the product does not act with water as calcium oxide does, and this lime is not adapted to the preparation of ordinary mortar. On the other hand, it gradually becomes solid, in contact with water, for reasons which are not known. Such substances are known as cements, or hydraulic cements. \ Calcium Sulphide, OaS, is formed by heating calcium sulphate with charcoal. It is converted by water«into the hyflrosulphide and the hydroxide : 2CaS + 2H 2 = Ca(SH) 2 + Ca(OH) 2 . The hydrosulphide is easily soluble in water, while the hydroxide is difficultly soluble. If the water is saturated with the hydroxide, it does not act upon the sulphide. It is remarkable on account of the fact that it is phosphores- cent. After having been exposed to sunlight, it continues to give light for some time afterward. This and the similar compound, barium sulphide, are now used in the preparation of luminous objects, such as match-boxes, clock-faces,* plates for house-numbers, etc. ^ Calcium Nitride, Ca 3 ]S[ 2 , a bronze-colored mass, is formed by heating calcium in nitrogen. It is decomposed 5*° COLLEGE CHEMISTRY. by water, with the formation of calcium hydroxide and ammonia. •^ Calcium Carbide, C 2 Ca, is formed by heating lime and carbon together in an electric furnace, when the reaction represented by the following equation takes place : CaO + 30 = C 2 Ca + 00. It is manufactured on the large scale, the form of carbon used being coke. The carbide forms a crystalline mass. That of average quality has a reddish color. That of bad quality has a grayish or black color. When pure it is transparent and white. When treated with water it yields acetylene, C 2 H 2 (which see). Strontium, Sr (At. Wt. 87.6) 1 *h Occurrence and Preparation — Strontium occurs in nature in the form of the sulphate, SrS0 4 , as celestite, and in the form of the carbonate, SrC0 3 , as strontianite. The latter is found in large quantities in Westphalia. The element is isolated by the action of an electric current on the molten chloride. Properties. — It is similar to calcium. It is oxidized by contact with the air, and decomposes water rapidly with evolution of hydrogen, which does not, however, take fire spontaneously. Compounds of Strontium. — The compounds of strontium are similar to those of calcium. Its chloride has not the same power to combine with water that calcium chloride has, though it deliquesces when left in contact with the air. The oxide is not easily made by decomposition of the carbonate by heat, as the carbonate is much more stable than that of calcium. It is, however, prepared without difficulty by heating the nitrate. When brought in con- tact with water, the oxide forms the hydroxide,. •which is analogous to calcium hydroxide. It is more easily soluble in water than the latter. Strontium nitrate, Sr(N0 3 ) 2 , is made in considerable BARIUM. 511 quantity for the purpose of preparing a mixture which, when burned, gives a red light (red-fire, Bengal-fire). It is easily made by dissolving strontianite or strontium car- bonate in nitric acid. Strontium sulphate. SrS0 4 , occurs in nature in beauti- ful crystals as the mineral celestite. It is formed when a soluble sulphate is added to a solution of a strontium salt. Its solubility lies between that of calcium sulphate and that of barium sulphate. Barium, Ba (At. Wt. 137.4). Occurrence and Preparation — Barium occurs in nature in the same forms of combination as strontium, viz., as the carbonate, BaOO,, in witherite; and as the sulphate, BaS0 4 , in barite or heavy spar. It is prepared by electrol- ysis of the molten chloride. Properties. — It closely resembles calcium and strontium, being a yellow metal, which is oxidized by contact with the air and readily decomposes water at the ordinary tem- perature. Barium Chloride, BaCl 2 + 2H a O, is prepared by dissolv- ing barium carbonate in hydrochloric acid. It dissolves easily in water, but not as easily as the chlorides of strontium and calcium. The order of solubility, beginning with the most soluble, is, calcium, strontium, barium, — the same as in the case of the sulphates. Barium Hydroxide, Ba(OH) a , is formed by dissolving barium oxide in water, just as calcium hydroxide is formed by treating calcium oxide with water. In hot water it is much more easily soluble than calcium hydroxide, and it is also more easily soluble in cold water. As such a solu- tion acts in the same general way as lime-water, it is sometimes used in the laboratory for the purpose of detecting carbon dioxide, barium carbonate being insolu- ble. Like lime-water, it has an alkaline reaction. Barium Oxide, BaO, is made by heating the nitrate, as the carbonate is not easily decomposed by heat. The most interesting property of the oxide is its power to take 5" COLLEGE CHEMISTRY. up oxygen when heated to dark red in the air or in oxygen, when it forms ^Barium Peroxide or Dioxide, Ba0 2 . — This peroxide is a white powder that looks like the simple oxide. When heated to a temperature a little higher than that required for its formation, it breaks down into barium oxide and oxygen. The formation of the peroxide by heating the oxide in the air, and the decomposition of the peroxide at a higher temperature, make it possible to extract oxygen from the air and to obtain it in the free state. This method of preparing oxygen on the large scale from the air was referred to under Oxygen. It is stated that the oxide improves with use. Specimens which have been in use for two years are said to be as efficient as at first. When a solution of hydrogen dioxide, H 2 2 , is added to a solution of barium hydroxide, a precipitate is formed which has the composition Ba0 2 -f- 8H 2 0. When filtered and put in a vacuum over sulphuric acid, it loses all its water and leaves behind pure dioxide. The dioxide is a con- venient starting-point in the preparation of hydrogen dioxide. It is only necessary to treat it with hydrochloric acid in order to make a solution of hydrogen dioxide. The solution made in this way, however, contains barium chloride. To make a solution containing nothing but the dioxide, pure barium peroxide is treated with dilute sul- phuric acid, when insoluble barium sulphate is formed and the hydrogen dioxide remains in solution: Ba0 2 + H 2 S0 4 = BaS0 4 + H 2 2 . It is interesting to compare the action of hydrochloric acid on barium peroxide and on the corresponding com- pound of manganese. As we have seen, with the latter the reaction takes place as represented in this equation : Mn0 2 + 4HC1 = Mn01 2 + 2H 2 + Cl 2 ; while with barium peroxide the reaction takes place thus: BaO, + 2HC1 = BaCl 2 + H 2 2 . B/tRlUM COMPOUNDS. 513 It is probable that in the case of manganese dioxide some intermediate reactions take place which are impossible in the other case. (See Manganese Dioxide.) Barium Sulphide, BaS, is made as calcium sulphide is, by reducing the sulphate by heating with charcoal. It is phosphorescent, like the calcium compound. b>- Barium Sulphate, BaS0 4 . — This occurs in nature as barite, or heavy spar, and is precipitated when a soluble sulphate or sulphuric acid is added to a solution of a barium salt. It is insoluble in water; when freshly pre- cipitated, it is easily soluble in concentrated sulphuric acid. It is artificially prepared for use as a pigment and is known as permanent white. On account of its insolu- bility it is much used in chemical analysis for the purpose of detecting and estimating sulphuric acid. \ Barium Carbonate, Ba0O 3 , occurs in nature as wither- ite, and is made pure by adding ammonium carbonate and a little ammonia to a solution of barium chloride. The carbonate usually found in the market is made by precipi- tating a solution of the crude sulphide with sodium car- bonate, or by heating together sodium carbonate and natural barium sulphate, or heavy spar. Made in either way it contains alkaline carbonate, from which it is im- possible to separate it by washing. The carbonate, like the other salts of barium, is poisonous. It has the power to unite, and form insoluble compounds, with metallic oxides of the formula M,0 3 , as, for example, ferric oxide, Pe,O s , and is used in analytical operations for the purpose of separating iron from other metals, like manganese, which are not precipitated by it. Phosphates of Barium. — The phosphates of barium correspond in general to those of calcium. Reactions which are of Special Value in Analysis. — The sulphates of calcium and strontium are completely converted into the carbonates by contact with a solution of ammonium carbonate in ammonia. The sulphate of barium is not changed in this way. Consequently, if a mixture of the three sulphates is treated with ammonium 514 COLLEGE CHEMISTRY. carbonate, those of calcium and strontium will be con- verted into carbonates, while that of barium will remain unchanged. By filtering, washing with water, and treat- ing with dilute nitric or hydrochloric acid, the carbonates will be dissolved, while the sulphate will not. If nitric acid is used, the solution may be evaporated to dryness, and treated with a mixture of alcohol and ether. Calcium nitrate will dissolve; strontium nitrate will not. Fluosilicic acid produces a precipitate of barium fluo- silicate, BaSiP 6 , in solutions of barium salts. This is insoluble in a mixture of alcohol and water, and difficultly soluble in water. The corresponding salts of calcium and \strontium are soluble. Calcium sulphat-e'in solution produces a precipitate in a solution of a strontium salt or of a barium salt, but not in one of a calcium salt. Strontium sulphate in solution precipitates barium sul- phate from a solution of a barium salt, but forms no pre- cipitate in a solution of a strontium salt. When boiled with a solution of one part of sodium carbonate and three parts of sodium sulphate, the sul- phates of strontium and calcium are completely convene I into carbonates, while the sulphate of barium remains unchanged. Barium chloride is insoluble in absolute alcohol; calcium chloride is easily soluble; while strontium chloride dis- solves in warm absolute alcohol. Ammonium oxalate, (NH 4 ) 2 C 2 4 , produces precipitates of the oxalates in solutions of calcium, barium, and stron- tium. Only the calcium salt is insoluble in dilute acetic acid. Potassium bichromate, K 2 Cr 2 7 , precipitates barium chromate, BaCr0 4 . The corresponding salts of calcium and strontium are soluble in water. Barium chromate is easily soluble in hydrochloric or nitric acid. All three elements of the group give colored flames which have characteristic spectra. Calcium compounds color the flame reddish yellow; strontium compounds give an color. MAGNESIUM. 5^ intense red; and barium compounds a yellowish green uoior. The spectra are more complicated than those of the elements of the potassium group, but each one contains highly characteristic lines which are easily recognized. MAGNESIUM SUB-GROUP. Glucinum, Gl (At. Wt. 9.1). Occurrence and Preparation. — The principal form in which the element glucinum occurs in nature is in the mineral beryl, which is a silicate of aluminium and glucinum of the formula Al a Gl 3 (Si0 3 ) 6 . Emerald has the same composition, but is colored green by the presence of a little chromic oxide. The element can be isolated by decomposing the chloride by heating it with potassium or sodium. Magnesium, Mg (At. Wt. 24.36). Occurrence. — Magnesium occurs abundantly in nature, though by no means as abundantly as calcium. Among the widely distributed minerals that contain the element are magnesite, which is the carbonate, MgC0 3 ; dolomite, a double carbonate of magnesium and calcium; serpentine, talc, soapstone, meerschaum, hornblende, all of which contain magnesium silicates. Further, the metal is found in solution in many spring-waters in the form of the sul- phate, or, as it is called, Epsom salt. Kainite is a sul- phate and chloride of the composition expressed by the formula K 2 S0 4 . MgS0 4 . MgCl 2 + 6H 2 0: kieserite is mag- nesium sulphate, MgS0 4 +H s 0; carnallite is a double chloride, KMgOl, + 6H 2 0. Magnesium compounds are contained in the soil in consequence of the decomposition of minerals containing it. It is to some extent taken up by the plants, and subse- quently into the animal body. It is found in the bones and in the blood in small quantities. Preparation. — The metal is made by the electrolysis of its chloride. In practice carnallite, MgKCl s + 6H 2 0, is 5 l6 COLLEGE CHEMISTRY. used. This is dehydrated and then melted in an iron crucihle. A carbon rod brought into the molten mass serves as the anode, while the crucible itself acts as the cathode. The metal magnesium separates at the cathode. Properties — It is a silver- white metal with a high lustre. In the air it changes only slowly, but it gradually becomes covered with a layer of the hydroxide. At ordinary tem- peratures magnesium does not decompose water; at 100° it decomposes it slowly. When heated above its melting- point in oxygen or in the air, it takes fire and burns with a bright flame, forming the white oxide. The light of the flame is very efficient in producing certain chemical changes, such as those involved in photography, when a permanent impression is made by the light upon a sensitive plate. It has also the power to cause hydrogen and chlorine to combine just as the sunlight and the elecbric light do. Applications. — The principal use to which magnesium is put is for the purpose of producing a bright light, as for photographing in spaces to which the sunlight does not have access, and for signaling. It is also used to some extent as an ingredient of materials employed in making fireworks. Compounds of Magnesium. — The compounds of mag- nesium present a general resemblance to those of glucinum. As the element is much more abundant in nature, its compounds have been studied more extensively. Its acid properties are somewhat weaker, and its basic properties stronger, than those of glucinum. Its hydroxide does not form salts with the hydroxides of potassium and sodium. On the other hand, its chloride is decomposed when its water solution is evaporated to dryness. The hydroxide is very slightly soluble in water, and this solution has a slightly alkaline reaction. \ Magnesium Chloride, MgCl 2 . — This salt, as has been stated, occurs in nature. It is easily formed by dissolving magnesium oxide or carbonate in hydrochloric acid. On evaporating at as low a temperature as possible, there M/tGNESIUM OXIDE— MAGNESIUM SULPHATE. 517 finally crystallizes out of the very concentrated solution a salt of the composition MgCl 2 + 6H 2 0, analogous to crystallized calcium chloride, OaCl 2 -f- 6H a O, and stron- tium chloride, Sr01 2 -f- 6H 2 0. When this crystallized salt is heated for the purpose of driving off the water, it is completely converted into the oxide MgCl 2 + H 2 = MgO + 2HC1. The chloride is a white, crystalline mass that de- liquesces in the air. At a bright red heat, it is volatile and can be distilled in an atmosphere of hydrogen. v ^\ Magnesium Oxide, MgO. — This compound is commonly called magnesia. A fine white variety, known as magnesia usta, is made by heating precipitated basic magnesium carbonate. It is a white, loose powder, Avhich is difficultly soluble in water, forming with it the hydroxide, Mg(OH) 2 , which is also difficultly soluble. Magnesia is used, in medicine, as an application to wounds, and, mixed with a solution of ferric sulphate, as an antidote in cases of poisoning by arsenic. As magnesia is infusible, it is used to protect vessels which are subjected to a high tempera- ture. When mixed with water and allowed to lie in the air, it becomes very hard. Mixtures of magnesia with sand also have this property, and are used as hydraulic cements. It is used, further, in the manufacture of fire- bricks. Magnesium Sulphate, MgS0 4 . — The mineral kieserite, which occurs in Stassfurt, has the composition expressed by the formula MgS0 4 + H 2 0; or, more probably, this should be written (HO) 2 MgS0 3 , in which it appears as a derivative of the acid SO(OH) 4 . The salt, MgS0 4 + 7H 2 (or H 2 MgS0 5 + 6H 2 0), also occurs in nature. It is this variety which is generally obtained when a solution of magnesium sulphate is evaporated to crystallization. It crystallizes in large rhombic prisms, or, if rapidly deposited from very concentrated solutions, in small, needle-shaped crystals. At ordinary temperatures, 100 parts of water 518 COLLEGE CHEMISTRY. dissolve 125 parts of the salt. The water solution has a bitter, salty taste. When heated, it readily loses 6 mole- cules of water, but it requires a temperature of over 200° to drive off the last molecule. This has led to the belief that the salt with one molecule has the constitution above given, being a derivative of the acid SO(OH) 4 . Applications. — Magnesium sulphate finds extensive ap- plication. It is used in medicine as a purgative, and is known as Epsom salt, for the reason that it is contained in the water of Epsom springs; it is used further in the manufacture of sodium sulphate and potassium sulphate, and as a fertilizer in place of gypsum, it having been shown to be advantageous in some cases. Its chief use is for loading cotton fabrics. ~^ Magnesium Carbonate, MgCO s . — Magnesium shows a marked tendency to form basic salts with carbonic acid. When a neutral magnesium salt is treated with a soluble carbonate, a basic carbonate is precipitated, the composi- tion of which varies according to the conditions under which it is prepared. The salt obtained by adding an excess of sodium carbonate to a solution of magnesium sulphate has the composition, Mg 3 (OH) 2 (C0 3 ) 2 . The basic carbonate manufactured on the large scale is known as magnesia alba. It is this form of the carbonate that is used in the preparation of magnesia usta. Normal mag- nesium carbonate, MgC0 3 , occurs in nature as magnesite. It crystallizes in the same form as calcium carbonate, or is isomorphous with it. It is insoluble in water, but like calcium carbonate it dissolves in water containing carbon dioxide in solution. From this solution crystals having the composition MgC0 3 + 3H 2 and MgCO s + 5H 2 are deposited under the proper conditions, r- Phosphates. — The conduct of the phosphates of mag- ,/ nesium is very similar to that of the phosphates of calcium. All three are known; and of these only the primary salt is soluble in water. A salt much utilized in analysis is ammonium-magnesium phosphate, Mg(NH 4 )P0 4 . This is difficultly soluble in water, and may therefore be used BORATES, SILICATES, ETC. 519 either for the purpose of detecting magnesium or phos- phoric acid. The corresponding salt of arsenic acid, Mg(NH 4 )As0 4 , is very similar to the phosphate, and on account of its insolubility it is also used in chemical analysis. p( Borates — A borate of magnesium together with mag- ) nesium chloride occurs in nature, and is known as boracite. It has the composition expressed by the formula 2Mg B + MgCl a . The borate, Mg 3 B 8 15 , is derived from the acid, H 6 B 8 16 , which is related to normal boric acid, as is shown by the equation 8B(OH) 3 = H 6 B 8 15 + 9H 2 0. - p Silicates.— The simplest silicate of magnesium found in /nature is olivine, which is represented by the formula * Mg s Si0 4 . It is the neutral -salt of normal silicic acid. Serpentine is derived from the acid, 0 aDd has the composition Mg 3 Si 2 7 -f 2H 2 0. y Magnesium Silicide, Mg,Si, is made by heating together magnesium chloride, sodium fluosilicate, sodium chloride, and sodium. Under these circumstances the sodium sets magnesium free from the chloride, and silicon from the fluosilicate. Both unite to form magnesium silicide. When treated with hydrochloric acid it gives . silicon hydride, SiH 4 , and hydrogen:- Mg 2 Si + 4H01 = 2Mg01 2 + SiH 4 . The liberation of hydrogen is due to the presence of an excess of magnesium. Reactions of Magnesium Salts which are of Special Value in Chemical Analysis. — Soluble hydroxides (KOH, NaOH, NH 4 OH) precipitate magnesium hydroxide. If ammonium chloride is present ammonia does not precipi- tate the hydroxide. Di-sodium phosphate with ammonia and ammonium chloride precipitates ammonium-magnesium phosphate from the solution of a magnesium salt. 520 COLLEGE CHEMISTRY. Sodium and potassium carbonates precipitate basic magnesium carbonate. Ekbium, E (At. Wt. 166). As regards the position of erbium in the periodic system, a final statement cannot as yet be made. It is even ques- tionable whether it is an element. EXPERIMENTS. Calcium Salts. Experiment 218. — Dissolve 10 to 20 grams of limestone or marble in common hydrochloric acid. Filter, and evaporate to dryness. Expose a few pieces of the residue to the air. Magnesium and its Salts. % Experiment 219. — Make anhydrous magnesium chloride thus : Dissolve 180 grams magnesia usta in ordinary hydrochloric acid; shake the solution with an excess of magnesia to remove iron and aluminium ; filter ; add 400 grams ammonium chloride ; evapo- rate to dryness, keeping the mass constantly stirred. The double salt thus formed must be dried until a small specimen put in a test-tube is found not to give off water when heated. The dry salt is then ignited in a crucible placed in a furnace until ammo- nium chloride is no longer given off, when the molten mass, which is anhydrous magnesium chloride, is poured out on a stone and, after it is broken up, it is put in a dry bottle provided with a good stopper. CHAPTEE XXVII. ELEMENTS OF FAMILY III, GROUP A: ALUMINIUM.— SCANDIUM.— YTTRIUM.— YTTERBIUM.— SAMARIUM.— HELIUM. "^ General. — There is in some respects a resemblance between boron and the principal member of this group; but as boron acts almost exclusively as an acid-forming element, it was treated of in connection with the elements of Family V, Group B, or the nitrogen group. Attention was, however, called to the fact that the analogy between these elements and boron is but slight. The points of resemblance between boron and the members of Family III, Group A, will be indicated below. The principal member of this group is aluminium. The others are all rare, aDd some have been but imperfectly studied, owing to serious difficulties in the way of obtaining their com- pounds in pure condition. They are trivalent in their compounds, the general formulas being such as the follow- ing: MC1 3 , M(OH) 3 , M(N0 3 ) 3 , M 3 (S0 4 ) 3 , M a (C0 3 ) 3 , MP0 4 , etc. Aluminium oxide is weakly basic, and somewhat acidic, though less so than boron oxide. Aluminium hydroxide has the power to neutralize most acids, and also to form salts with strong bases. Boron oxide, on the other hand, has scarcely any basic properties, though it does form a few extremely stable compounds, in which the boron takes the place of the hydrogen of acids. (See Boron Phos- phate, p. 342.) 521 52 2 COLLEGE CHEMISTRY. Aluminium, Al (At. Wt. 27.1). \ Occurrence. — Aluminium is an extremely important element in nature and in the arts. It occurs widely dis- tributed, and abundantly in many different forms of com- bination. Among them are feldspar, mica, cryolite, bauxite. Feldspar is a silicate of aluminium and potas- sium of the formula AlKSi 3 8 . Mica is a general name applied to a large number of minerals which are silicates of aluminium and some other metal, as potassium, lithium, magnesium, etc. The simplest form of mica is that repre- sented by the general formula KAlSi0 4 , according to which the mineral is a salt of orthosilicic acid, Si(OH) 4 . Cryolite is a double fluoride of aluminium and sodium, or the sodium salt of fluoaluminic acid, Na s AlF 6 . Bauxite is a hydroxide of aluminium in combination with a hydroxide of iron. Aluminium also occurs in the products of decom- position of minerals. One of the most important of these is clay, which is found in all conditions of purity — from the white kaolin to ordinary dark-colored clay. Kaolin is the aluminium salt of orthosilicic acid of the formula Al 4 (Si0 4 ) 3 + 4H a O. Aluminium silicate is found in all soils, but is not taken up by plants, and does not find entrance into the animal body. The name aluminium has its origin in the fact that the salt alum was known at an early date, and the metal was afterwards isolated from it. x Preparation. — The preparation of aluminium on the large scale presents a problem of the highest importance. The element has properties which adapt it to many uses to which iron is put, and for many purposes it has advantages, over iron. Further, we are supplied by nature with un- limited quantities of the compounds of aluminium, which are distributed everywhere over the earth. While, how- ever, iron, lead, tin, copper, and other metals can be isolated from their natural compounds without serious difficulty, aluminium, which is more abundant than any of them, and in many respects more valuable than any of them, is locked in its compounds so firmly that it is only ALUMINIUM. S23 by comparatively complicated and expensive methods that it can be isolated. The best method for the preparation of aluminium con- sists in the electrolysis of aluminium oxide, in the form of corundum. The oxide is contained in iron crucibles, and carbon rods are introduced into the molten mass. The iron crucible forms one electrode and the carbon rods the other. \ Properties. — The color of aluminium is like that of tin, and it has a high lustre. It is very strong, and yet malleable. It is lighter than most metals in common use, its specific gravity being 2.5 to 2.7 according to the condition, while that of iron is 7.8, that of silver 10.57, that of tin 7.3, and that of lead 11.37. It does not change in dry or in moist air; and in the compact form it does not act upon water even at elevated temperatures. It melts at about 700°, which is higher than the melting- point of zinc, and lower than that of silver. Hydrochloric acid dissolves it with ease, forming aluminium chloride. At the ordinary temperatures nitric and sulphuric acids do not act upon it; at higher temperatures, however, action takes place, and the corresponding salts are formed. It dissolves in solutions of the caustic alkalies, forming the so-called aluminates. It reduces many oxides when heated with them to a sufficiently high temperature; and is used in the preparation of boron and silicon. When a mixture of finely divided aluminium and ferric oxide is heated by means of a burning magnesium wire or by means of a cartridge made of magnesium powder and barium peroxide, the aluminium is oxidized at the expense of the ferric oxide, and an extremely high temperature is produced. This is the brightest white heat (about 3000°) which is sufficient to melt the iron and also the aluminic oxide which are formed. A method of welding has been devised by G-oldschmidt which is based upon this reaction. C^ Applications. — The metal is used to a considerable extent in the preparation of ornaments, and of useful articles in which lightness is a matter of importance, as 5^4 COLLEGE CHEMISTRY. » in telescopes, opera-glasses, cooking-utensils, etc. An alloy with a small percentage of silver is used for the beams of chemical balances. Aluminium bronze, which is an alloy with copper, is also used quite extensively. It will be again referred to under Copper. ^ Aluminium Chloride, A1C1 3 . — When aluminium hydrox- ide is dissolved in hydrochloric acid a solution of alu- minium chloride is formed, and from this solution a compound of the formula A1C1 S + 6H 2 can be obtained in crystallized form. Like calcium and magnesium chlo- rides, this salt is deliquescent. When heated to drive off the water the salt conducts itself like magnesium chloride, but the decomposition into the oxide and hydrochloric acid takes place more easily than that of magnesium chlo- ride. The reaction is represented by the equation 2A1C] 3 + 3H 2 = A1 S 3 + 6H01. The dry chloride is prepared by the same method as that used in the preparation of silicon chloride and boron chloride, viz., by passing chlorine over a heated mixture of the oxide and carbon. The chloride, being volatile, sublimes, and is deposited in the cool part of the vessel, when pure, as a white, laminated crystalline mass. Generally, however, it is more or less colored in conse- quence of the presence of impurities. When exposed to the air it attracts moisture and gives off hydrochloric acid. It dissolves in water very easily, with a marked evolution of heat, but, from what was said above, it is evident that it cannot be obtained from this solution again by evapora- tion. It is volatile without change. The specific gravity of its vapor has been determined by different observers, and, unfortunately, with different results. According to Deville and Troost, it is such as to lead to the formula A1 2 C1 6 . Nilson and Pettersson, however, have found it to correspond to that required by the formula A1C1 S , their determinations having been made at a higher temperature than those of Deville and Troost. Still later determina- tions by Crafts again lead to the formula A1 2 C1 6 . In view ALUMINIUM HYDROXIDE— ALUMINATES. 525 of the conflicting evidence, the formula for aluminium chloride used here is the simpler one. By means of it and similar formulas for the other compounds of aluminium, the reactions of the element can be expressed somewhat . more easily and probably just as truthfully as by means of the more complicated formula. \ Aluminium Hydroxide, Al(OH),.— Normal aluminium ydroxide, Al(OH),, occurs in nature as the mineral hydrargillite. It is precipitated from a solution of aluminium chloride by ammonia : A1C1 3 + 3NH 4 OH = Al(OH), + 3NH 4 C1. Obtained by precipitation it forms a gelatinous mass, which is suggestive of starch-paste, and it is on this account extremely difficult to wash it completely free from the substances in the solution. It dries in the air, form- ing a gummy substance which has the composition Al(OH),. When heated under proper conditions it loses water, and forms the compound A10 a H : Al(OH) s = AIO.OH + H a O. This compound is found in nature as the mineral diaspore. If heated to a higher temperature it is converted into the oxide, A1 2 0, : 2A1(0H) 3 = ALO, + 3H 3 0. In the conduct of the chloride and of the hydroxide aluminium exhibits a certain resemblance to boron. The acidic character of the latter is, however, more strongly marked than that of the former. Boron chloride is more easily decomposed by water than aluminium chloride, and, as the decomposition takes place at the ordinary tem- perature, the product is the hydroxide instead of the oxide, as in the case of aluminium. The hydroxide, B(OH) 3 , readily loses water and forms metaboric acid, which in composition is analogous to diaspore; and at a . higher temperature the oxide, B 2 0, , is formed. V Aluminates. — When sodium or potassium hydroxide is added to a solution of an aluminium salt, aluminium 526 COLLEGE CHEMISTRY. hydroxide is at first precipitated, but an excess of the reagent dissolves the precipitate. This action is due to the acidic character of aluminium hydroxide. It is prob- able that in solution the action with potassium and sodium hydroxides is of the same kind as represented in the equa- tions Al(OH) 3 + 3KOH = Al(OK) 3 + 3H 2 0, and Al(OH), + 3NaOH = Al(ONa), + 3H„0. On evaporating the solution of the potassium salt, how- ever, the product obtained has the formula AIO.OK, and is plainly the salt of the hydroxide AlO. OH, which may be called nieta-aluminic acid, to suggest its analogy to metaboric acid, BO. OH. When aluminium hydroxide and sodium carbonate are melted together, the salt AlO.ONa is formed, as has been shown by determining the amount of carbon dioxide given off when a known weight of the hydroxide is employed. When, however, the solution of the hydroxide in caustic soda is evaporated, the salt Al(ONa) s is deposited. These salts are v,ery unstable, though their solutions can be boiled without undergoing decomposition. Carbon dioxide decomposes them at once with precipitation of aluminium hydroxide. Similar salts are formed with calcium and barium. Among them may be mentioned those of the following formulas: Ca s ( A10 s ) 2 , Ca(A10,) a , Ba s (A10 3 ) 2 , and Ba(A10 2 ) 2 . The calcium salts are insolu- ble in water, and some of them become hard in contact with water. They are therefore of importance in the manufacture of hydraulic cements. The barium salts are soluble in water. Many aluminates occur in nature, forming the impor- tant group of minerals known as the spinels. Of these, spinel itself is the magnesium salt of the hydroxide AlO. OH, and is represented by the formula ,,a'q>Mj, or Mg(A10 2 ) 2 . Chrysoberyl is the corresponding glucinum salt G1(A10 2 ) 2 ; and gahnite is the zinc salt Zn(A10 a ) 2 . ALUMINA TES. 527 These salts are extremely stable, differing markedly in this respect from those above referred to, which are made in the laboratory. They are decomposed by heating them, in finely powdered condition, with primary or acid potas- sium sulphate, the action of which was described on p. 466. As will be seen farther on, there are other salts similar to the aluminates in structure which occur in nature. Among these chromic iron, or chromite, which is an iron salt of a hydroxide of chromium of the formula CrO.OH, may be mentioned here. The salt is to be regarded as made up according to the formula ^ ^ ^ > Fe, or Fe(CrO,),. Further, magnetic oxide of iron or mag- netite, Fe 3 4 , is regarded as belonging to the same group, and its constitution represented thus: ^ 6 QQ>'Fe, or Fe(FeO,) s ; and there is also a compound of magnesium, -p, e rt ' >Mg. For the sake of emphasizing these anal- ogies, the formulas of the compounds above mentioned are here presented in tabular form : Potassium aluminate AIO.OK Sodium aluminate AlO. ONa Calcium aluminate .1^ ' q> Oa Barium aluminate »i q" q > Ba Spmel A10.0 >Mg Chrysoberyll. Alo!o >G1 Gahnite . jq Chromite q v q ' q> Fe Magnetite -p q Magnesio-ferrite FeO > ^» 528 COLLEGE CHEMISTRY. \ Aluminium Oxide, Al a O s . — As has been stated, the oxide is formed by heating the hydroxide. It is found in nature in the form of ruby, sapphire, and corundum. The natural variety is extremely hard; and granular corun- dum, which is known as emery, is used for polishing. The red color of the ruby is caused by the presence of a trace of a chromium compound; while the blue color of the sapphire is probably due to the presence of a trace of a cobalt compound. Aluminium oxide is infusible in the hottest furnace fire, but it melts in the flame of the oxy- hydrogen blowpipe, and on cooling it becomes crystalline. In the electric furnace it melts and is vaporized. By mixing it with various easily fusible substances and heat- ing, it is obtained in the form of crystals, and by adding certain metallic oxides these crystals can be colored. In this way artificial rubies and sapphires have been prepared, which have all the properties of the natural ones. When the oxide is moistened with a few drops of a solution of cobaltous nitrate and then ignited, it turns blue. This fact is taken advantage of in chemical analysis for the purpose of detecting aluminium. When the oxide is made by gently igniting the hydroxide, it dissolves in strong acids. If, however, it is heated to a high tempera- ture, acids will not dissolve it. The natural varieties of the oxide, further, are not soluble in acids. By fusion with acid potassium sulphate insoluble aluminium oxide is converted into a soluble compound. ~" Aluminium Sulphate, A1 2 (S0 4 ) 3 . — This salt is made by dissolving the hydroxide of aluminium in dilute sulphuric acid, and evaporating to crystallization, when a salt of the composition A1 2 (S0 4 ), + 18H 2 is deposited. When heated the salt loses its water of crystallization, and, if the temperature is raised to that of red heat, the anhydrous salt is decomposed with loss of sulphur trioxide and forma- tion of iluminium oxide. This decomposition is, however, *iot complete. The sulphate is manufactured on the large scale for various purposes, as, for example, for a mordant, for sizing paper> etc. ALUMS. 529 Alums — When a solution of aluminium sulphate is brought together with a solution of potassium sulphate in the proportion of their molecular weights, a salt crystal- lizes out which has the composition represented by the formula KA1(S0 4 ) 2 + i2H 2 or K 2 S0 4 + A1 2 (S0 4 ), + 24H 2 0. The most rational view held regarding this compound is SO ^ 0K that it has the constitution . &r\ ^ O-pAl, with perhaps some of the so-called water of crystallization present in the form of hydroxyl. This salt, which has long been known under the name of alum, is the type of a class of similar compounds, all of which are called alums. These may be regarded as derived from the ordinary form by substituting sodium, ammonium, or any other member of the sodium group, besides some other metals, for the potassium. Thus a series of alums is obtained, of which the following are examples : NaAl(S0 4 ), + 12H 2 LiAl(SG 4 ) 2 + 12H 2 (NH 4 )A1(S0 4 ) 2 + 12H a O CsAl(S0 4 ) 2 + 12H 2 0. Again, alums are derived from the ordinary form by sub- stituting for the aluminium some other elements that resemble aluminium, as, for example, iron, chromium, and manganese. Such alums are those represented by the following formulas : KFe(S0 4 ) 2 + 12H 2 0; KCr(S0 4 ) 2 +12H 2 0; KMn(S0 4 ) 2 + 12H 2 0. All the alums have certain properties in common. They are all soluble in water, and all crystallize in the same forms, which are regular octahedrons combined with 53° COLLEGE CHEMISTRY. cubes. If a crystal of one alum is suspended in the solu- tion of any other one it will continue to grow. They are all strictly isomorphous. The principal alums containing aluminium are those of potassium and ammonium, both \of which are manufactured on the large scale. Potassium Alum, Potassium - Aluminium Sulphate, KAl(SOJ 2 + ISHjO.— Ordinary alum is found in nature in some volcanic regions. The mineral alunite, which is a basic salt of the formula K(A10) 3 (S0 4 )., + 3H a O, or per- haps K[Al(OH),] 3 (S0 4 ) a , occurs in larger quantities. When this salt is heated and treated with water, ordinary alum dissolves, and is easily obtained from the solution. Another source of alum is alum shale. This occurs in large quantities in nature, and consists of coal, clay, and iron pyrites. When it is heated in contact with the air the coal burns, as do also the sulphur and pyrites, and sulphuric acid is formed. When allowed to lie for a time in contact with the air the iron pyrites is converted into sulphate and sulphuric acid. The latter acts upon the clay or aluminium silicate, forming aluminium sulphate, from which alum can easily be made. It is easier to treat the shale and similar substances with sulphuric "acid directly, and this method is now generally employed. Alum dissolves readily in hot water, 357.5 parts of the crystallized salt dissolving in 100 parts of water at 100°. At 0° only 3.9 parts dissolve, and at the ordinary tempera- ture about 12 parts. It crystallizes beautifully in regular octahedrons, occasionally with cube faces developed on them. Under some circumstances it crystallizes in cubes. When heated, alum melts in its water of crystallization, and if heated to a sufficiently high temperature the water passes off, leaving burnt alum. Applications. — Alum is used very extensively in the preparation of pigments, as a mordant, in the sizing of paper, for clarifying water, etc. Ammonium Alum, Ammonium-Aluminium Sulphate, (NH 4 ) A1(S0 4 ) S + 12H 2 0, is in most respects much like the potassium compound, and can be used in pla".e of it for ALUMimUM SILICATES. \ 531 almost all purposes for which alum is used. It is some- what more easily soluble than ordinary alum. As it is cheaper than the latter it is largely manufactured in place of it. A Sodium Alum is much more easily soluble in water than /either potassium or ammonium alum, and this makes it difficult to prepare it in pure condition. It is therefore not manufactured, although sodium compounds are cheaper than those of either potassium or ammonium. Aluminium Silicates. — It has been stated that alumin- ium silicate enters into the composition of a number of important minerals. It occurs in enormous quantities in nature. The most important of the minerals containing it are the feldspars, of which ordinary feldspar, KAlSi 3 8 , and albite, or sodium feldspar, NaAlSi,O e , are the most abundant. These again enter into the composition of granite together with quartz and mica, and mica is -itself a double silicate of aluminium. As remarked under Silicic Acid (which see), the natural silicates are for the most part salts of polysilicic acids which are derived from ortho- silicic acid by loss of water from two or more molecules. Up to the present but little more has been done with the many natural silicates than to determine their percentage composition. It appears probable from what has already been learned regarding their constitution that investiga- tions in this direction will before long yield interesting results. As yet, however, the methods for such investiga- tions are quite unsatisfactory, owing largely to the fact that the compounds are so extremely stable that but few reagents decompose them, and if they are decomposed at all, the products are such that no conclusion can be drawn regarding the constitution. A careful study of the rela- tions in which minerals occur in nature will undoubtedly be of assistance, as this will throw some light upon the conditions under which they were formed. One of the most common decompositions of minerals, constantly taking place, and that has taken place to an enormous extent, is that of feldspar. Under the influence 532 COLLEGE CHEMISTRY. of moisture and the carbon dioxide of the air, this sub- stance slowly decomposes, the products being mainly potassium or sodium silicate and aluminium silicate. The salt of the alkali metal, principally potassium, being solu- ble, is carried away, and finds its way into the soil. The silicate of aluminium is not soluble, but it easily forms emulsions with water, and is therefore carried down the sides of the hills and mountains upon which it is formed into the valleys, and much of it finds its way into streams, Sometimes this carrying away is prevented, and then large beds of comparatively pure clay, known as kaolin, are formed. The clay found in the valleys is always more or less impure and colored. Kaolin. — This is the purest form of aluminium silicate found in nature. It always contains water. Its composi- tion varies, some specimens on analysis giving results which lead to the formula Al 4 (Si0 4 ) 3 + 4H 2 0, according to which the substance is the salt of normal silicic acid, Si(OH) 4 . Other specimens have the composition HAlSi0 4 fOH + H,0, or Si -j q\ a i I tt Q. When heated alone kaolin does not melt; but if feldspar is added to it, the whole melts, and forms a translucent mass known as porcelain. Other substances besides feldspar may be used for this purpose. Clay. — Clay, as has been stated, is a name given to tne impure varieties of aluminium ^silicate which have been carried down from the place of formation. Among the substances besides aluminium silicate found in clays are calcium carbonate, magnesium carbonate, sand, and hydroxides of iron. The color is largely determined by the amount of the hydroxides of iron present. The better varieties are used in the manufacture of the so-called " stone -ware," gas-retorts, and fire-bricks. The colored varieties are used for making ordinary earthenware and bricks. Marl is clay mixed with considerable quantities of calcium carbonate. PORCELAIN. 533 Ultramarine. — The substance occurring in nature and known as lapis lazuli consists of a silicate of sodium and aluminium together with a sulphur compound, probably a polysulphide of sodium. The coloring matter, known as ultramarine, obtained by powdering it was formerly very expensive, but it is now made artificially by the ton, and the color of the artificially prepared substance is even more beautiful than that of the natural. ^ Porcelain. — It was stated above that when kaolin is * heated alone it does not melt, but that when feldspar is added to it, or when that found in nature contains feld- spar, as is frequently the case, it either fuses, forming a compact mass, or it melts and forms a translucent mass. Further, when kaolin or any other variety of clay is mixed with water, a plastic substance results, that can be kneaded and worked into any desired form. These facts form the basis of the manufacture of earthenware, porcelain, etc. The ease with which the mass melts depends upon the quantity of feldspar or other flux added to it. If but little is added it melts with difficulty; if much is added it melts easily. In the manufacture of the finest kinds of porcelain kaolin is used. This is generally mixed with a little feld- spar or chalk, gypsum or some other flux, and sand is also added. All these substances must be very finely ground. The mixture is then worked into the desired forms, and carefully dried. After the objects are dried they are next burned, first at a red heat at which -the mass becomes solid, afterwards at a white heat for the purpose of form- ing a glaze upon the surface. The product after the first burning is that which is familiar as porous earthenware ; that formed in the second burning is the porcelain with glaze as it is commonly used. In order to form the glaze upon the porcelain the porous earthenware first formed is drawn through a vessel con- taining proper materials in finely powdered condition and suspended in water. The materials used are generally the same as those used for the porcelain itself, but they are 534 COLLEGE CHEMISTRY. mixed in different proportions, with less kaolin, and more sand and feldspar, so as to make them more easily fusible. After this treatment the objects are again heated to a high temperature. Earthenware. — The ordinary varieties of earthenware are made from varieties of clay which are much less pure than kaolin. Ordinary colored clay is used. The objects are formed, and then subjected in general to the same kind oi treatment as porcelain. They are glazed in different ways. One method consists in bringing the glazing material on the earthenware before it is burned; another method con- sists in putting the objects in the furnace without a glaze, and towards the end of the firing process sodium chloride is thrown into the furnace, and is thus brought in contact with the ware in the form of vapor. A chemical change takes place, resulting in the formation of a silicate of aluminium and sodium upon the surface. This melts, and forms a glaze. Bricks are the most common variety of unglazed earthenware. Owing to the presence of other substances besides aluminium silicate, as, for example, calcium car- bonate, the material is comparatively easily fusible. The color of bricks is largely due to the presence of oxides of „ iron. ^ Alloys. — Aluminium forms alloys with magnesium. Those containing 10 to 25 per cent magnesium come into the market under the name of magnalium. This has a specific gravity 2 to 2.5, and is therefore lighter than aluminium, but it is much harder. Magnalium containing a considerably higher percentage of aluminium is used in making mirrors, for which it is well adapted on account of its hardness and its power of taking a high polish. Aluminium amalgam is an extremely active substance. It decomposes water energetically. Reactions of Aluminium Salts which are of Special. Value in Chemical Analysis. — Potassium- and sodium hydroxides precipitate aluminium hydroxide, which is soluble in an excess of the reagents. SC/tNDlUM. 535 Ammonia precipitates the hydroxide, which is only slightly soluble in an excess of the reagent. Hydrogen sulphide and carbon dioxide precipitate aluminium hydroxide from a solution of an aluminate; that is, from a solution of aluminium hydroxide in a caustic alkali. Ammonium sulphide and other soluble sulphides precipi- tate the hydroxide. This is due to the instability of the sulphide of aluminium, or, going farther back, to the weak basic character of the hydroxide. The reaction of ammonium sulphide with aluminium sulphate takes place as represented in the following equation : Al,(SO0> + 3(NH 4 ) a S + 6H a O = 3(NH.) a S04 + 3H S S + 2Al(OH),. Soluble carbonates precipitate aluminium hydroxide for the same reason that the soluble sulphides do. The reaction between aluminium sulphate and sodium car- bonate takes place thus : AlaflSOO. + 3Na 2 OO s + 3H a O = 3Na,S0 4 + 3CO a + 2Al(OH) s . Other Members of Family III, Group A. Scandium, Sc (At. Wt. 44.1). — This element was dis- covered in 1880 in the minerals euxenite and gadolinite. Its compounds are similar to those of aluminium. It forms an oxide of the formula Sc a 3 ; a sulphate, Sc 2 (S0 4 ) 3 ; a double sulphate, KSc(S0 4 ) 3 ; etc. It is of special interest for the reason that its properties were foretold by Men- deleeff several years before it was discovered. The prophecy was based upon the position of the element in the periodic system. When the relations between the atomic weights and properties of the elements were first described in a comprehensive way by Lothar Meyer and Mendeleeff, the latter described the properties of an ele- ment then unknown, which he called ekaboron. This should have the atomic weight about 44, should form an oxide of the formula M,0 3 , etc. It has been shown that the properties of scandium agree very closely with those foretold. 53 6 COLLEGE CHEMISTRY. Yttrium, Y (At. Wt. 89), like scandium, is found in gadolinite, euxenite, and some other rare minerals. The element itself has not been isolated. Ytterbium, Yb (At. Wt. 173).— This rare element, like scandium and yttrium, is found in gadolinite and euxenite — most abundantly in the latter. Its compounds in general resemble those of yttrium. Samarium, Sm, has been obtained from samarskite, a North Carolina mineral. It has also been found in small quantities in some other minerals, as cerite, gadolinite, and orthite. In general it resembles aluminium in its chemical conduct. ^^ Helium. — When cleveite and some other rare minerals ^ are treated with acids n gas is given off. This was at first held to be nitrogen, but a careful spectroscopic examina- tion by Kamsay has shown that it contains a hitherto unknown gas. . The spectrum contains lines which are identical with lines that have been observed in the solar spectrum. These have been ascribed to an unknown element, helium. This has now been isolated and studied. It proves to be a very light gas, as incapable of forming chemical compounds as argon. Practically all the evidence goes to show that it is an element, but as yet little is known in regard to its relations to other elements. The chemistry of lanthanum is so intimately connected with that of cerium and didymium, that, although these three elements appear to belong to different families, they will be briefly treated of together. EXPERIMENT. Aluminium Chloride. Experiment 220.— Aluminium chloride is made thus : Mix aluminium oxide with starch-paste ; form the mass into small balls the size of ordinary marbles ; ignite these in a crucible in a furnace ; put them in a porcelain tube, and then pass dry chlo- rine over them, at the same time heating the tube to redness. The chloride will sublime and be deposited in the front end of the tube or in a receiver if the heat is sufficient. It can be purified by subliming it over heated iron or aluminium. CHAPTEE XXVIII. ELEMENTS OF FAMILY I, GROUP B: COPPER— SILVER.— GOLD. General. — The facts which strike one most forcibly on comparing the elements of this group with those of Group A of the same family are, that they are much less active chemically, and that they furnish a greater variety of compounds. Sodium and potassium and the other members of Group A display the greatest activity, as we have seen. The basic character is most strongly developed in them. Further, in nearly all their compounds they act with the same valence. They are univalent in all their salts. Copper, silver, and gold, however, are not chemic- ally active elements, and the activity grows less with increasing atomic weight. Copper and gold form two series of compounds each, and silver also forms a few com- pounds, in which it appears with a valence greater than one. In the two series of salts formed by copper tb element appears to be univalent and bivalent, as in tl v chlorides CuCl and CuCl 2 . Gold, however, -is univalent and trivalent, while silver is almost exclusively univalent. It must be said that the resemblance between gold and the other members of Group B, Family I, is apparently not as marked as that between mercury and copper and silver. Copper, Cu (At. Wt. 63.6). General. — The compounds of copper which are most commonly met with are those in which it acts as a bivalent element. Its principal compounds are copper oxide, CuO; 537 538 COLLEGE CHEMISTRY. the sulphate, CuS0 4 ; and the sulphide, CuS. In all these the copper is bivalent. But besides these there are such compounds as CuCl and Cu a O, in which the element appears to be univalent. There are, then, two series of salts, of which the following will serve as examples : CuCl CuCl, CuBr CuBr, Cu,0 CuO Those compounds which are of the first order, corre- sponding to the chloride CuCl, are called cuprous com- pounds. Thus, CuCl is cuprous chloride; Cu a O, cuprous oxide, etc. On the other hand, compounds of the second order are called cupric compounds. Thus, CuCl a is cupric chloride; CuO, cupric oxide; CuS0 4 , cupric sulphate, etc. Forms in which Copper occurs in Nature. — Copper is a widely distributed element, and it occurs also in large quantities. It occurs in the uncombined condition, or as native copper, in large quantity in the United States in the neighborhood of Lake Superior, in China, Japan, Siberia, and Sweden. The most valuable ores of copper are the oxides, ruby copper or cuprous oxide, Cu a O, and cupric oxide, CuO; the carbonates, as malachite, Cu a (OH) 2 CO s ; the sulphides, as chalcocite, Ou a S; copper pyrites, Cu a S.Fe a S s ; and others. Metallurgy of Copper — The metallurgy of copper is comparatively complicated, owing to the difficulty of con- verting the ores of copper into the oxide. In most of the ores used sulphur and iron are contained, as well as smaller quantities of other elements, as arsenic, antimony, lead, etc. The ores are first roasted with the object of converting the sulphides partly into oxides. Under these circumstances the sulphides of iron are more easily con- verted into the oxides than the sulphides of copper. By adding a material rich in silicic acid, and melting the roasted ore in a blast furnace with charcoal, the oxide of iron is partly reduced, and converted into silicate, which runs off with the slag. In this way a product is obtained COPPER. 539 which is richer in copper than the roasted ore. This, which is called the matte, contains copper sulphide and iron sulphide. The matte is again roasted and melted in the same way as the ore, and a further quantity of iron is removed, while some of the copper is reduced. A reaction which plays an important part in these processes is that which takes place between cuprous oxide and cuprous sul- phide, forming metallic copper and sulphur dioxide : 2Cu 2 + Cu„S = 6Cu + SO,. Sometimes it is necessary to repeat the roasting and melt- ing with charcoal and sand a number of times, the matte becoming richer in copper at each successive stage. Properties. — Copper is a hard metal, of a reddish color and metallic lustre. It does not change in dry air, but in moist air it gradually becomes covered with a green layer of a basic carbonate. It melts at a somewhat lower tempera- ture than gold, and at a somewhat higher temperature than silver. It is very malleable and tenacious. It decomposes water only at bright-red heat. When heated in the air to a comparatively high temperature it becomes covered with a layer of cupric oxide; at a lower temperature cuprous oxide is formed. Nitric acid dissolves it, copper ni- trate, Cu(NO,),, being formed, and the oxides of nitrogen being evolved (see p. 286); dilute sulphuric acid does not act upon it unless the air has access to it; concen- trated sulphuric acid when heated with it forms cupric sulphate, CuS0 4 , and sulphur dioxide. Dilute acids in general do not act upon it unless the air has access to it. This fact is of importance in connection with the use of copper vessels in culinary operations. Substances con- taining vegetable acids can be boiled in bright copper vessels with impunity, for the water-vapor prevents the access of air, but, on cooling, the air is admitted, and then action takes place, causing solution of some of the copper, which is objectionable. Ammonia in contact with copper absorbs oxygen, and the copper dissolves in consequence of the formation of a compound of cupric oxide and am- 540 COLLEGE CHEMISTRY. monia. This fact is sometimes taken advantage of for the preparation of nitrogen, as has already been stated (see p. 251). Applications. — As is well known, copper is used very extensively for a variety of purposes, among which the following may be mentioned : for electrical apparatus, coins, copper vessels, roofs, for covering the bottoms of ships, etc. It is also used for copper-plating; and in the preparation of a number of valuable alloys, such as brass, bronze, gun-metal, bell-metal, etc. Alloys. — Brass is a mixture or compound of about one part of zinc and two parts of copper; these proportions may, however, be varied between quite wide limits. There is a variety of brass containing equal parts of zinc and copper, and another containing one part of zinc and five parts of copper. Pinchbeck is made by combining two parts of copper and one of brass. Bronze consists of copper, zinc, and tin. The propor- tion of copper varies from 65 to 84 per cent; that of zinc from 11 to 31.5 per cent; and that of tin from 2.5 to 4 per cent. When exposed to the air bronze becomes covered with a green coating of basic copper carbonate, which protects it from further action. This coating is now generally produced artificially by a variety of methods, as by washing the surface with a solution of salts and acids. Gun-metal consists generally of copper and tin in the proportion of 11 parts of tin and 100 parts of copper. Bell-metal contains a larger proportion (from 20 to 25 per cent) of tin than gun-metal. Alloys with Aluminium containing aluminium and copper in widely different proportions are made. That with 3 per cent of copper has a whiter color than alumin- ium, the color being more like that of silver. On the other "hand, an alloy of copper with 5 to 10 per cent of aluminium has a color resembling that of gold. This, which is known as aluminium bronze, is very hard and elastic, and is not easily acted upon by chemical reagents. COMPOUNDS OF COPPER. 541 It is now used to a considerable extent in the manufacture of ornamental and useful articles. German silver is an alloy consisting of copper, zinc, and nickel. The proportion of copper varies from 40 to 60 per cent; that of zinc from 19 to 44 per cent; and that of nickel from 6 to 22 per cent. Cuprous Chloride, CuGl, is formed by heating cupric chloride, CuCl,; and by boiling an excess of copper filings with concentrated hydrochloric acid with the addition of a little nitric acid, filtering through asbestos, and pouring into water. It is a white crystalline compound, and is difficultly soluble in water. When exposed to the air it turns green in consequence of the formation of a basic chloride. It is volatile at a high temperature, and a determination of the specific gravity of the vapor gave a result corresponding to the " formula Cu.Cl,. It has markedly the power to combine with chlorine, and there- fore acts as a reducing agent. Cupric Chloride, CuCl 2 .— This compound is formed by treating copper or cuprous chloride with chlorine. It is also easily made by dissolving cupric hydroxide, or car- bonate, in hydrochloric acid. From its solution in water the chloride crystallizes with two molecules of water, CuCl a + 2H 2 0. The crystals when heated lose their water without suffering further decomposition, except at high heat, when a part of the chlorine is given off, and cuprous chloride is formed. Cupric chloride combines with am- monia gas, forming a compound of the formula CuCl 2 . 6NH 3 , which is soluble in water, with a dark blue color. When heated it loses four molecules of ammonia, and the com- pound CuCl 2 .2NH 3 is left behind. Cuprous Hydroxide, Ou(OH). — The simple compound of the formula here given is not known, but a derivative of this, of the formula Cu 8 3 (OH) 2 or 40u J O.H a O, is easily made by adding sodium hydroxide to a solution of a cuprous salt. It passes readily over into cuprous oxide when gently heated. Cuprous Oxide, Cu 2 0, occurs in nature, and is known 542 COLLEGE CHEMISTRY. as ruby copper or cuprite. It is easily prepared by treating a solution of glucose, or starch sugar, with copper sulphate and potassium hydroxide. By boiling, the copper is thrown down in the form of cuprous oxide. At first this is yellow, and it is supposed by some that the yellow com- pound is the hydroxide, but satisfactory evidence of this has not been furnished. The yellow precipitate is soon converted into the red oxide. Cuprous oxide is not changed when allowed to lie in contact with the air. It dissolves in nitric and sulphuric acids, forming cupric salts; and if the acids are dilute, copper is deposited. This will be clear from a consideration of the following equation : Cu 2 + H 2 S0 4 = CuS0 4 + H 2 + Cu. Cupric Hydroxide, Cu(OH) 2 , like the hydroxides of most base-forming elements, is thrown down by the addi- tion of a soluble hydroxide to a cupric salt. It is a voluminous, blue precipitate. When allowed to stand in a solution, or when the solution is boiled, the hydroxide loses a part of its hydroxyl, and is converted into a black compound of the formula Ou(OH) 2 -\- 2CuO, and this when dried and heated is converted into the oxide CuO. > Cupric Oxide, CuO. — Cupric oxide is found in nature ' iu the neighborhood of Lake Superior in the United States, •md is formed by heating copper to redness in contact with the air, or by heating the nitrate. It loses its oxygen very readily when treated with reducing agents, such as hydrogen and carbon. It is used extensively in quantita- tive analysis for the purpose of estimating the composition of organic compounds, or such as contain carbon and hydrogen. Its use is based upon the fact that when organic compounds are heated with the oxide they are oxidized, the carbon being converted into carbon dioxide and the hydrogen into water. By passing the products of the oxidation through calcium chloride, and a solution of potassium hydroxide, the water is retained in the first, CUPRIC SULPHATE. 543 and the carbon dioxide in the second, and the weight of each formed can easily be determined. Cupric Sulphate, CuS0 4 . — This salt is manufactured on the large scale, and in the crystallized form, containing five molecules of water, CuS0 4 + 5H 2 0, is commonly called "blue vitriol" or "blue stone." It is found to some extent in nature, being formed by the action of the oxygen of the air on the sulphide. It is most conveniently made by dissolving metallic copper in concentrated sulphuric acid, or by treating cupric sulphide with sulphuric acid. The action of sulphuric acid on the metal consists essen- tially in the formation of cupric sulphate, sulphur dioxide, and water, as expressed in the equation Ou + 2H 2 S0 4 = CuS0 4 + SO, + 2H„0. ' The question whether the copper reduces the sulphuric acid directly, or the hydrogen given off from the acid effects the reduction, is an open one. But there are other products formed besides those mentioned. At first a brown substance of the composition Cu 2 S is deposited. As the action proceeds oxysulphides are formed, the final product of a series of changes being Cu 2 OS, or CuO.CuS, which is black and insoluble in water. Under some con- ditions a considerable proportion of the copper is trans- formed into the oxysulphides by sulphuric acid. Cupric sulphate is obtained in large blue crystals of the triclinic system, which have the composition CuS0 4 -f- 5H 2 0. When heated to 100°, four molecules of water are given off, and the last is not given off until the temperature 200° is reached. This makes it appear probable that the salt has the constitution represented by the formula CuSO s (OH) 2 , corresponding in this respect to magnesium sulphate (which see). When heated higher, it loses all its hydroxyl, and the salt, CuS0 4 , is left in the form of a white powder, which has the power to take up water from the air, becoming blue again. It dissolves in three parts of cold water and one-half part boiling water. Copper 544 COLLEGE CHEMISTRY. sulphate, containing seven molecules of water, CuSO, -j- 7H 3 0, is obtained when mixed with solutions of the sul- phates of iron, zinc, or magnesium, all of which crystallize with seven molecules of water. In this form cupric sul- phate is isomorphous with the other sulphates. These salts have in general received the name of vitriols, and the old names "green vitriol," "white vitriol," and "blue vitriol " are still used to some extent, though rarely by chemists. Among the similar salts included under the same general head are the following : Zinc sulphate (white vitriol) ZnS0 4 -f 7H„0 Magnesium sulphate MgS0 4 + 7H 2 Glucinum sulphate G1S0 4 + 7H..0 Ferrods sulphate (green vitriol) ]?eS0 4 + 7H a Nickel sulphate NiS0 4 + 7H S Cobalt sulphate CoS0 4 -f 7H a Copper sulphate (blue vitriol) CuS0 4 + 7H 2 (CuS0 4 + 5H 2 0) Applications. — Cupric sulphate is used extensively in the preparation of blue and green pigments, in copper- plating by electrolysis, in galvanic batteries, for the pur- pose of preserving wood, and as a remedy against phylloxera (see p. 382), etc. Cupric Nitrate, Cu(N0 8 ) 2 , is easily formed by dissolving copper in dilute nitric acid. It is easily soluble in water, and is deposited in crystallized form, the crystals contain- ing three or six molecules of water according to the tem- perature, the salt with six molecules being formed at the lower temperature. Like other copper salts, it has a blue color. Cupric Arsenite, CuHAs0 3 , is formed as a greenish- yellow precipitate when an ammoniacal solution of arseni- ous acid is added to a solution of cupric sulphate. It is known as Scheele's green. A compound of cupric arsenite and cupric acetate, which is made by treating a basic acetate of copper with arsenious acid, is known as Paris green or Schweinfurt green. CUPRIC CARBONATES, ETC. 545 Cupric Carbonates. — When a soluble carbonate is added to a solution of cupric sulphate a voluminous greenish precipitate is formed, which has the composition Cu< Cu(OH) a .CuC0 3 , or q>CO. This is plainly a basic Cu< 0H carbonate. The mineral malachite, which has a beautiful green color, has the same composition as the precipitate just mentioned. Cuprous Sulphide, Cu 2 S.— This compound occurs in nature, and is known as chalcocite. It is, further, a con- stituent of copper pyrites, which is a compound of cuprous and ferric sulphides, CiijS.Fe^, or OuFeS 2 . It can be made by heating copper and sulphur together in the right proportions. It has a grayish-black color; does not give up its sulphur, even when heated in hydrogen; and is the more stable of the two sulphides of copper. Cupric Sulphide, OuS. — This is formed as a black pre> cipitate when hydrogen sulphide is passed into a solution of a cupric salt. In water alone cupric sulphide is some- what soluble. Hence in washing out a precipitate of copper sulphide with water a little of it will pass through in solution. It also easily undergoes oxidation, and, as it forms the sulphate, some is dissolved in this way unless proper precautions are taken. It is slightly soluble in ammonium sulphide, but insoluble in sodium sulphide. The above facts are of importance in the analysis of com- pounds containing copper, as will readily be seen. When heated, cupric sulphide loses half its sulphur, and is con- verted into cuprous sulphide. Copper-plating. — The process of copper-plating consists in brief in depositing upon an object a layer of copper by putting it in a bath containing some copper salt, and connecting it with one pole of an electric battery. Decom- position of the copper salt takes place, and copper is deposited upon the object. Alkaline solutions of the double cyanides are best adapted to the purpose. The 546 COLLEGE CHEMISTRY. process is extensively used in the preparation of electrotype plates. These are prepared either from wood-cuts or from type by making a mould of gutta-percha, covering this with graphite, and immersing the plate thus prepared in the copper-plating bath. The plate thus made is an exact reproduction of the wood-cut or type of which the impres- sion in gutta-percha was taken. Reactions which are of Special Value in Chemical Analysis. — Potassium or sodium hydroxide forms a blue precipitate which becomes black on standing or when heated. (See Oupric Hydroxide.) Ammonia first forms a greenish precipitate, which is a basic salt. With cupric sulphate the reaction takes place thus: SO,<2>Cu Cu<£ H " +2NH,+3H 1 0= ">SO,+ (]raj,S0 4 . S0 2 Cu Cu Co. Cadmium, Cd (At. Wt. 112.4). General. — The compounds of cadmium are similar to those of zinc and magnesium. The element occurs in nature in much smaller quantity than either of these, CADMIUM. 565 often in company with zinc, and its compounds are not as frequently met with. It is always bivalent. A mineral known as greenockite is cadmium sulphide, OdS. Preparation and Properties. — Cadmium is obtained principally from different varieties of zinc blende, and separates with the zinc. Being more volatile than zinc, it passes over first when the mixture is distilled. From this first distillate, which contains the oxides of zinc and cadmium, the metals are reduced by heating with char- coal. It has a color like that of tin, and is harder than tin. According to the specific gravity of its vapor, its molecule is identical with its atom, for the molecular weight is approximately 112. Cadmium chloride, CdCl 2 , like zinc chloride, is volatile ; the sulphate crystallizes well, but is not analogous in com- position to the sulphates of magnesium and zinc, as the composition of the crystallized salt is represented by the formula 3CdS0 4 + 8H Q 0; the normal carbonate, CdCO s , is precipitated by soluble carbonates. Cadmium Sulphide, CdS, is one of the most character- istic compounds of the element. It is a beautiful yellow substance, that is thrown down from a solution of a cad- mium salt by hydrogen sulphide. While it dissolves in concentrated acids it does not dissolve in dilute acids, and it is therefore completely precipitated by hydrogen sul- phide. It is used as a constituent of yellow paints. Analytical Reactions. — Cadmium, as has just been stated, is precipitated by hydrogen sulphide. It is thrown down together with the other elements of the hydrogen sulphide group (see p. 214). As the sulphide is not solu- ble in ammonium sulphide, it is easily separated from those of arsenic, antimony, and tin by treating with this reagent, when it is left undissolved in company with the sulphides of mercury, lead, bismuth, and copper. The doable salt of cuprous cyanide and potassium cyanide is not decomposed by hydrogen sulphide, whereas the corre- sponding salt of cadmium is decomposed by it, and the yellow sulphide is precipitated. 5 66 COLLEGE CHEMISTRY. The hydroxide of cadmium .differs from that of zinc in not having acid properties. It does not dissolve in the caustic alkalies. Mercury, Hg (At. Wt. 200.3). General. — As already stated, mercury yields two series of compounds, known as mercurous and mercuric com- pounds, which are analogous to the two series of copper compounds. While, however, copper forms with the oxygen acids such salts as belong to the cupric series, as CuS0 4 , Cu(N0 3 ) 2 , etc., mercury forms salts belonging to both series. There is, for example, a mercurous nitrate, HgNO s , and a mercuric nitrate, Hg(NO s ) 2 ; a mercurous sulphate, Hg 3 S0 4 , and a mercuric sulphate, HgS0 4 , etc. The mercurous compounds are readily converted into the mercuric compounds by the action of oxidizing agents, and the mercuric are converted into mercurous compounds by the action of reducing agents. The action will be treated of under the individual compounds. The question as to the correct formula of the mercurous salts is in the same condition as that in regard to the formula of cuprous salts, with this difference : the molecular weight of mer- curous chloride leads to the formula HgOl, but there is evidence that when the chloride is heated some mercury is set free, and this has led to the suggestion that the mole- cule corresponds to the formula Hg 3 Cl 2 , and that the compound breaks down into mercury and mercuric chlo- ride when heated. It is, however, quite possible that the compound has the simpler formula, and that this under the influence of heat is partly decomposed, as represented in the equation 2HgCl = HgCl 2 + Hg. The fact that mercury is set free is, therefore, by no means satisfactory evidence that the formula of mercurous chloride is Hg,Cl 2 , and in the present state of the inquiry it is justifiable to write the formula HgCl. Forms in which Mercury occurs in Nature. — Mercury occurs native to some extent, but principally in the form MERCURY. 567 of the sulphide, HgS, which is known as cinnabar. This is sometimes found crystallized, but generally amorphous. The chief localities are Idria, Almaden in Spain, and New Almaden in California. Metallurgy of Mercury. — In order to obtain mercury from the sulphide this is roasted in vessels so constructed as to condense and collect the vapor of mercury given off. In the roasting process the sulphur is oxidized to sulphur dioxide, which of course escapes. In some places the ore is mixed with limestone and distilled from clay or iron retorts, when the mercury passes over. Crude mercury is redistilled in order to purify it. It is also purified by treating it with dilute nitric acid or with a solution of ferric chloride. Properties. — Mercury is a silver-white metal of a high lustre. At ordinary temperatures it is liquid, though at — 39.5° it becomes solid. Its specific gravity is 13.5959. It does not change in the air at ordinary temperatures. It boils at 357.25°, and is converted into a colorless vapor, the specific gravity of which leads to the conclusion that, as in the case of cadmium, the molecule and atom are identical, or that the molecule consists of only one atom. It is insoluble in hydrochloric acid and in cold sulphuric acid; but dissolves in hot concentrated sulphuric acid, and is easily soluble in nitric acid. The vapor of mercury is very poisonous. Applications. — Mercury is extensively used in the manu- facture of thermometers, barometers, etc. ; as tin-amalgam for mirrors; and m the processes by which gold and silver are obtained from their ores. Amalgams. — The alloys which mercury forms with other metals are called amalgams. These compounds are generally obtained without difficulty simply by bringing mercury in contact with other metals. Among the amal- gams which are of chief interest are those of sodium, ammonium, silver, and gold. Sodium amalgam is made by bringing mercury and sodium together. A crystallized amalgam containing the constituents in the proportions 568 COLLEGE CHEMISTRY. represented in the formula Hg 6 Na has been obtained. Generally, sodium amalgam is easily decomposed by water, the mercury separating in the free state and the sodium acting upon the water, forming hydrogen and sodium hydroxide. It is much used in the laboratory as a con- venient means of producing hydrogen in alkaline solutions. It serves as an excellent reducing agent in some cases. Ammonium amalgam has already been referred to under the head of Ammonia (which see). It is a curious sub- stance, which is formed when an electric current acts upon a solution of ammonia containing some mercury which is connected with the negative pole, and also very easily by pouring a solution of ammonium chloride upon sodium amalgam. In the latter case sodium chloride and am- monium amalgam are formed. Apparently the reaction takes place according to the following equation : NH 4 C1 + NaHg = NaCl + NH.Hg. The product is extremely voluminous, and swells up during the reaction, so that it occupies under favorable circumstances about twenty times the volume occupied by the sodium amalgam. It has a metallic lustre, resembling in general the other amalgams. It is very unstable at the ordinary temperature, breaking down into mercury, hy- drogen, and ammonia. At a low temperature, however, it has been obtained in crystallized form. The metallic lustre and general outward appearance of the compound suggests that whatever is in combination with mercury in it has probably metallic properties, and this affords some confirmation of the ammonium theory, according to which the presence of the complex, NH 4 , in the salts formed by ammonia is assumed. Silver amalgam and gold amalgam vary in composition according to the method of prepara- tion, and when heated are comparatively easily decomposed. Mercurous Chloride, HgCl, is commonly called calomel. Like cuprous chloride, CuCl, and argentic chloride, AgCl, it is insoluble in water. It is formed most readily by reducing mercuric chloride. The reduction can be accom- MERCUROUS CHLORIDE— MERCURIC CHLORIDE. 569 plished by means of sulphurous acid, when the following reaction takes place : 2HgCl 2 + 2H 2 + SO a = 2HgCl + H 2 S0 4 + 2H0L It is also formed by heating together mercuric chloride and mercury, and by subliming a mixture of mercuric sulphate, sodium chloride, and mercury. This method is the one mostly used in the manufacture of calomel. The product obtained by sublimation is crystalline ; the precipi- tated substance forms a loose powder. As was stated above, the specific gravity of the vapor corresponds to that required for the formula HgCl. When acted upon for some time by light it undergoes partial decomposition into mercury and mercuric chloride. This is a fact of great importance, inasmuch as calomel is much used in medi- cine, and mercuric chloride is an active poison. Bottles in which calomel is kept should be carefully protected from the action of the light. Just as mercuric chloride is converted into mercurous chloride by reducing agents, so the latter is converted into the former by oxidizing agents. When, for example, mercurous chloride is treated with nitric acid it is con- verted into mercuric chloride and mercuric nitrate. Mercuric Chloride, or Corrosive Sublimate, HgCl 2 , is made by subliming a mixture of sodium chloride and mer- curic sulphate, HgS0 4 + 2NaCl = HgCl 2 + Na 2 S0 4 , and by dissolving mercury in aqua regia, evaporating to dryness, and subliming the residue. It is a white, trans- parent, crystalline mass; it is soluble in water, and can be obtained in crystalline form from the solution. It is more easily soluble in alcohol and ether than in water, and is extracted from a water solution by shaking with ether. It is quite volatile, and the specific gravity of its vapor corresponds to that required for the formula HgCl 2 . It is easily reduced to mercurous chloride by contact with organic substances, and by reducing agents in general. 57° COLLEGE CHEMISTRY. The action of sulphur dioxide has already heen treated of as furnishing a method for the preparation of mercurous chloride. Stannous chloride abstracts chlorine from it and forms mercurous chloride and metallic mercury, while the stannous chloride is converted into stannic chloride: 2HgCl a + SnCl 2 = 2HgCl + SnCl 4 ; HgCl a + SnCl a = Hg + SnCl 4 . Mercuric chloride is an active poison, and has been used extensively in this capacity. It has a very marked influ- ence upon the lower organisms, which play such an im- portant part in producing disease and the decay of organic substances, and is used as a disinfectant. Wood impreg- nated with a solution of it is partly protected from decay. In surgery it is used for the purpose of preventing contamination of wounds by the hands and instruments of the surgeon, it being customary now for the surgeon to wash his hands and instruments in a dilute solution of the chloride before performing an operation. Mercurous Iodide, Hgl, can be made by treating mer- cury with iodine; or by treating mercuric iodide with mercury; and more easily by adding potassium iodide to a solution of a mercurous salt, when it is thrown down as a greenish-yellow powder. It is unstable, and breaks down, yielding mercury and mercuric iodide. When treated with potassium iodide it suffers the same decomposition, the iodide which is formed combining with the potassium iodide, while mercury is deposited. Mercurous iodide is used in medicine. As it is more easily decomposed than mercurous chloride, and mercuric iodide is poisonous, care should be taken in its use. Mercuric Iodide, Hgl 2 , is made by direct combination of mercury and iodine, and by the action of potassium iodide upon a solution of a mercuric salt, when it is pre- cipitated as a beautiful scarlet-red powder. Though insoluble in water, it is soluble in alcohol and ether. It dissolves, further, in a solution of potassium iodide in consequence of the formation of a double iodide, or iodo- MERCURIC IODIDE- MERCURIC OXIDE. 57 1 mercurate. When heated to about 150° the color changes from red to yellow, and after cooling the yellow substance changes to red. Sometimes it can be kept in the condition in which it has a yellow color for some time after it has cooled down, and then if touched at one point the change to the red substance takes place rapidly through the entire mass. Both the red and the yellow modifications are crystallized, but in different forms. The red crystals are tetragonal; the yellow ones are rhombic or monoclinic. The monoclinic structure is evidently, from what has been- said, an unstable one for this compound. It seems prob- able that the difference between the two forms is due to a difference in the complexity of the molecules; perhaps the larger molecules of the red iodide are decomposed by heat into the smaller molecules of the yellow iodide. When the iodide is first precipitated from a solution of mercuric chloride by potassium iodide it is yellow, but it rapidly turns red. Mercuric Oxide, HgO, like the iodide, presents itself in two colors — the red and the yellow. When mercury is . heated for a long time at a temperature near the boiling- point so that the air has access to it, it is converted into the red oxide. When a solution of a mercuric salt is treated with caustic soda or caustic potash the yellow oxide is precipitated. When heated, the red oxide becomes darker, and finally nearly black, and on cooling the red color reappears. The yellow oxide slso becomes darker on heating. Exposed to the light, the red oxide loses some of its oxygen, and mercury is deposited. When heated to a sufficiently high temperature both varieties give up all their oxygen, as was seen in the case of the red oxide in the preparation of oxygen. It will be remembered that oxygen was discovered by heating this compound. Of what importance this discovery was the student will now better appreciate than when the discovery was first men- tioned. It will now be seen that most of the chemical phenomena with which we have to deal involve the action of oxygen. S7 2 COLLEGE CHEMISTRY. Mercuric Sulphide, HgS, is the principal compound of mercury found in nature. This natural variety, which is known as cinnabar, is a red crystallized compound. When prepared by treating a solution of a mercuric salt with hydrogen sulphide it is an amorphous black powder. This black powder can, however, be converted into the red crystallized variety by sublimation, and by allowing it to stand in contact with a solution of the sulphide of an alkali metal. The sulphide, whether amorphous or crys- tallized, is acted upon with difficulty by acids. Dilute nitric acid, which easily dissolves the sulphides of those metals which resemble mercury, leaves it unacted upon, and advantage is taken of this fact for the purpose of separating it from the sulphides of lead, bismuth, copper, and cadmium in qualitative analysis. It is dissolved by concentrated nitric acid and by aqua regia. The substance used in medicine under the name Ethiops martialis is an intimate mixture of amorphous sulphide and sulphur. Like the oxide, cinnabar turns dark when heated, and it slowly undergoes the same change when exposed to the light, in consequence of a slight decomposition into mer- cury and sulphur. When that which has not been heated to too high a temperature is cooled down again, it acquires its original red color. If it has been heated to the tem- perature of sublimation, the black color of that which does not sublime is permanent. Cinnabar is used as a pigment. Mercuric Cyanide, Hg(CN) 2 , is formed by dissolving mercuric oxide in an aqueous solution of hydrocyanic acid. It is soluble in water, and crystallizes well in quadratic prisms. When heated it is decomposed into cyanogen and mercury, and this affords the most convenient method of making cyanogen. Fulminating mercury is an explosive compound much used in the manufacture of gun-caps. It is made by dis- solving mercury in nitric acid and adding alcohol. Its explosion consists in a sudden breaking down into nitro- gen, carbon dioxide, and mercury. The composition of the compound is represented by the formula C a N a 2 Hg. AMMONIA COMPOUNDS OF MERCURY. 573 Mercurous Nitrate, HgNO, , is made by treating nitric acid with an excess of mercury. This salt is easily decom- posed, forming difficultly soluble basic salts, some of which are of complicated composition. The simplest is that which has the composition HgOH.HgN0 3 . Mercuric Nitrate, Iig(NO s ) 2 , is formed by treating mercuric oxide with an excess of nitric acid and evaporat- ing, when the salt can, under favorable circumstances, be obtained in crystals. It is easily decomposed by water, with formation of a basic salt which is insoluble in water. Compounds formed by Salts of Mercury with Ammonia. — When mercuric chloride is treated with ammonia a white precipitate is formed, the composition of which is repre- sented by the formula Hg01NH 2 . The formation takes place according to the equation Chester ]■ HUH HgCl 2 + 2NH 3 = HgClNH, + NH 4 C1. The compound is known as white precipitate, or, to distin- guish it from another similar compound, as infusible white precipitate. The latter is formed by adding a solution of mercuric chloride to a boiling solution of ammonium chloride containing ammonia, and separates out on cooling. It has the composition Hg(NH 3 01) 2 . This is called mer- curic diammonium chloride, or fusible white precipitate. Mercurous compounds corresponding to both the above- mentioned derivatives of mercuric chloride are known. The first, or mercurous chlor amide, HgjNHjCl, is the black substance formed when mercurous chloride is treated with ammonia: 2HgCl + 2NH 3 = Hg 2 NH 3 Cl + NH^Ol. The second, or mercurous ammonium chloride, HgKH 3 Cl, is formed by treating calomel with ammonia gas. Similar compounds are obtained from the other salts of mercury. In a solution of mercurous nitrate ammonia forms a black precipitate, known in pharmacy as Mercurius solubilis Hahnemanni. Reactions which are of Special Value in Chemical Analysis. — As has been seen in the account already given 574 COLLEGE CHEMISTRY. of the conduct of the compounds of mercury, mercurous and mercuric compounds conduct themselves quite differ- ently. Among the most characteristic reactions of mer- curous compounds are the following : Sodium or potassium hydroxide gives a black precipitate of mercurous oxide. Ammonia g'ives a black precipitate which is a compound of mercurous oxide and ammonia. Hydrochloric acid and soluble chlorides form mercurous chloride, and the precipitate turns black when treated with ammonia. The reactions with stannous chloride, potassium iodide, hydrogen sulphide, and ammonium sulphide have been explained. The principal reactions of the mercuric salts are the following: Sodium or potassium hydroxide produces a yellow precipitate of mercuric oxide. The reactions with hydrogen sulphide, stannous chlo- ride, and potassium iodide have been described above. ELEMENTS OF FAMILY III, GROUP B: GALLIUM.— INDIUM.— THALLIUM. General. — All the elements of this group are rare. Gallium forms compounds which in composition are analogous to those of aluminium, in which it is trivalent, as in the compounds GaCl 3 , Ga 2 3 , Ga(NO s ) 3 , Ga,(S0 4 ) s , etc. On the other hand, it also forms compounds in which it is bivalent, as in the chloride GaCl, and GaO. Indium also forms compounds in which it is trivalent, and a few in which it is bivalent. Thallium, like gold, is tri- valent and univalent in its compounds. Gallium, Ga (At. Wt. 70). — This element is found in some varieties of zinc blende, and was discovered in that which occurs at Pierrefitte, in France. It owes its name to the Latin name of France, Gallia. Like scandium, it is of special interest for the reason that its properties were foretold by Mendeleeff some years before it was discovered, and it was described by him under the name of eka- INDIUM— THALLIUM. 575 aluminium, as scandium was described under the name of eka-boron. Indium, In (At. Wt. 114). — Indium was discovered in a variety of zinc blende found at Freiberg, in Germany. The discovery was made by means of the spectroscope, and as the spectrum of the element contains two very characteristic blue lines it was called indium. It has since been found in other varieties of zinc blende, but always in small quantity. Thallium, Tl (At. Wt. 204.1).— This element was dis- covered in the flue-dust of a sulphuric-acid factory in the Harz mountains, by the aid of the spectroscope. As it colors the flame a beautiful green it was called thallium, from the Greek ftaWos, which signifies a green branch. It has since been found in a number of varieties of iron pyrites and copper pyrites. EXPERIMENTS. Zinc and its Salts. \? Experiment 224. — Heat a small piece of zinc on charcoal in the oxidizing flame of the blowpipe. The white fumes of zine oxide (philosopher's wool) will be seen, and the charcoal will be covered with a film which is yellow while hot, but becomes white on cooling. Experiment 225. — Dissolve some zinc dust in a solution of sodium hydroxide, and see whether hydrogen is given off. Mercuet and its Salts. Experiment 226. — Make a solution of mercurous nitrate by treating at the ordinary temperature an excess of mercury with nitric acid, which is not too concentrated ; and with this solution study the conduct of mercurous salts. Experiment 227. — Heat some of the solution of mercurous 'iiitrate to boiling, then add a few drops of concentrated nitric acid, and boil again. With the solution thus obtained study the conduct of mercuric salts. CHAPTER XXX. ELEMENTS OF FAMILY IV, GROUP B: GERMANI UM.— TIN— LEAD. General. — Of the elements of this group germanium is extremely rare. It was discovered by Winkler in 1885. Tin and lead, on the other hand, have long been known, and are extensively used. All form two series of com- pounds, in one of which they are bivalent and in the other quadrivalent. The general formulas of some of the prin- cipal compounds of the first series are as follows : MO],, MO, M(OH)„ M(NO s ) 2 , MSO„ MCO„ etc. The general formulas of some of the principal' compounds of the other series are as follows : MC1 4 , M0 2 , M(OH) 4 , MO(OH) 2 , etc. These elements have already been referred to at the close of Chapter XXII (see p. 399), and attention was there called to the resemblance between them and carbon and silicon. In this connection it will be well to repeat what was there said. Of the three elements of the group, germanium and tin are more acidic in character than lead. They combine with chlorine in two proportions, forming the chlorides GeCl 2 , SnCl 2 , PbCl 2 , GeCl 4 , SnCl,, PbCl 4 . With oxygen they combine to form the compounds GeO,, SnO s , and Pb0 2 . Stannic oxide, Sn0 2 , and lead peroxide, Pb0 2 , form salts with bases, and these have the composi- tion represented by the general formulas M 2 SnO, and M 2 PbO s , and are therefore analogous to the silicates, car- 576 GERMANIUM— TIN. 577 bonates, and titanates. On the other hand, further, salts are known which are derived from the oxide PbO. These have the general formula M a Pb0 2 , and are to be regarded, as salts of an acid Pb(OH) 2 . These salts are not stable, and are not easily obtained. Most of the derivatives of lead are those in which it plays the part of a base-forming element. Notwithstanding the marked analogy between some of the compounds of tin and those of the members of the silicon group, it appears, on the whole, advisable to treat of this element in company with lead, which it also resembles in many respects. Germanium, Ge (At. Wt. 72). Germanium is the third element the properties of which were foretold by Mendeleeff by the aid of the periodic law. As it occurs in the silicon group he called it eka-silicon. It was discovered in a silver ore occurring at Freiberg, Germany. The name has, of course, reference to the country in which it was discovered. It acts mostly as a base-forming element, being perhaps more like tin than any other one metal. It forms the two chlorides GeCl a and GeCl 4 , and the corresponding fluorides GeP a and GeF 4 ; but preferably it forms those compounds in which the element is quadrivalent. The fluoride forms double salts resembling the fluosilicates, as, for example, potas- sium fluogermanate, K a GeP 6 . Tin, Sn (At. Wt. 118.5). General. — The compounds of tin with which we gen- erally have to deal belong, with the exception of stannous chloride, to the series in which the metal is quadrivalent, and in this series it acts as an acid-forming element. The chloride, SnCl 4 , corresponds to the chlorides of carbon and silicon, CC1 4 and SiCl 4 . Unlike these elements, how- ever, it does not form a compound with hydrogen. Occurrence. — Tin occurs almost exclusively as tin-stone or cassiterite in nature. This is the dioxide, Sn0 2 , corre- 578 COLLEGE CHEMISTRY. sponding to carbon dioxide, CO,; silicon dioxide, SiO,; titanium dioxide, TiO,; etc. It also occurs in small quantities in company with gold as metallic tin, and in a variety of pyrites of the formula Cu 4 SnS 4 + FeSnS 4 , known as stannite. Metallurgy. — The ores are roasted for the purpose of getting rid of the sulphur and arsenic, and the oxide is then heated with coal in a furnace. After the reduction is complete the tin is drawn off and cast in bars. This tin is impure, and when again slowly melted, that which first melts is purer. By letting it run ofE as soon as it melts, the comparatively difficultly fusible alloy remains behind, and the tin is thus rendered much purer. The commercial variety of tin known as Banco, tin is the purest. It receives its name from Banca, in the East Indies, where it is made. Block-tin is made in England, and is also comparatively pure. Properties. — Tin is a white metal, which in general appearance resembles silver. It is soft and malleable, and can be hammered out into very thin sheets, forming the well-known tin-foil. Its specific gravity is 7. 3. At 200° it is brittle, and at 228° it melts. At ordinary tempera- tures it remains unchanged in the air. It dissolves in hydrochloric acid, forming stannous chloride, SnCl,; in sulphuric acid, forming stannous sulphate, SnS0 4 , sulphur dioxide being evolved at the same time. Ordinary con- centrated nitric acid oxidizes it, the product being a compound of tin, oxygen, and hydrogen, known as meta- stannic acid, which is a white powder insoluble in nitric acid and in water. It is dissolved by a hot solution of potassium hydroxide which forms potassium stannate, K a SnO a . Applications. — It is used in making alloys, of which bronze (see p. 540), soft solder, and britannia metal are the most important. It is used also for protecting other metals, as in the tinware vessels in such common use, which are made of iron covered with a layer of tin. Copper vessels are also frequentlv covered with tin. STANNOUS CHLORIDE. 579 Alloys. — Bronze has already been treated of under Copper. It is made of copper, tin, and zinc. Soft solder is made of equal parts of tin and lead, or of two parts of tin and one of lead. Britannia metal is composed of nine parts of tin and one of antimony. Tin amalgam is made by bringing tin and mercury together, and is used in the silvering of mirrors. Alloys of tin and antimony are much used for machine bearings. Those in most common use contain not only these two elements, but also copper and Stannous Chloride, SnCl 2 , is formed by dissolving tin in hydrochloric acid, and if the solution is concentrated enough the compound crystallizes out. The crystals have the composition SnCl, + 2H 2 0. This is the commercial product known as tin salt. It is easily soluble in water, but if the solution is dilute it becomes turbid in conse- quence of the formation of the insoluble basic salt, 2Sn<^ H +H a O: 2SnC! [, + 3H 2 = 2Sn<° H + H 2 + 2HC1. This same decomposition takes place when the solution is allowed to stand in contact with the air. Under these circumstances a part of the stannous chloride is converted into stannic chloride by oxidation : 3SnCl 2 + H 2 + = SnCl 4 + 2Sn(OH)01. When the crystals are heated they melt at about 40°, an if heated higher they undergo partial decomposition, forming the oxide and hydrochloric acid : SnCl, + H 2 = SnO +2HC1. Stannous chloride has a marked tendency to combine with chlorine and to pass over into stannic chloride. This power has already been shown in its action upon mercuric chloride and upon mercurous chloride. It reduces the 580 COLLEGE CHEMISTRY. former, first to mercurous chloride, and it then abstracts the chlorine from this, leaving metallic mercury : 2HgCl a + SnCl a = SHgCl + SnCl 4 ; 2HgCl + SnCl a = 2Hg + SnCl 4 . Stannous chloride is a good mordant. Stannic Chloride, SnCl 4 , formed by treating tin with chlorine, is a colorless liquid which boils at 114°, and the specific gravity of its vapor is that required by a compound of the formula SnCl 4 . In the air it gives off fumes in consequence of the action of moisture. It has long been known by the name spiritus fumans Libavii. It combines with water, forming a number of crystallized hydrates. When its solution in water is boiled stannic acid is pre- cipitated : SnCl 4 + 3H 2 = H a Sn0 3 + 4HC1. Stannic chloride is used as a mordant. "With other chlorides it forms the chlorostannates of the general formula M a SnCl 6 or SnCl 4 .2M01, of which the ammonium salt, (NH 4 ) 2 SnCl 6 , or pink salt, is the best known. This is manufactured for use in making cotton -prints. Stannic Hydroxide, Sn(OH) 4 , is perhaps formed when a solution of stannic chloride in water is boiled. The precipitate obtained has, however, the composition H 2 Sn0 3 , and this is known as stannic acid. Stannic acid is precipi- tated also by treating a solution of a stannate with just enough acid to. effect decomposition. The decomposition with hydrochloric acid takes place as represented in the equation Na a Sn0 3 + 2H01 = 2NaCl + H a Sn0 3 . The compound thus obtained is insoluble in water, but is easily soluble in hydrochloric, nitric, and sulphuric acids, and in the caustic alkalies. With the alkalies it forms stannates, as sodium stannate, Na 2 Sn0 3 , and potassium STANNIC OXIDE, ETC. 581 stannate, K 2 SiiO s . The former is made on the large scale, and is known as preparing-salt. Metastannic Acid. — When tin is treated with concen- trated nitric acid it is converted into a white powder which is insoluble in water and in acids, but nevertheless seems to be a hydroxide of tin of the same composition as stannic acid. This is known as metastannic acid. With alkalies it forms salts which in properties and composition are entirely different from the stannates. They are known as the metastannates. Two sodium salts are known, which differ in composition. One of these has the composition Np,Sn 5 O u , the other is ]STa s Sn 9 19 . They are probably derived from acids analogous to the polysilicic acids, which may be called polystannic acids. The question as to the composition of metastannic acid is an open one. Stannic Oxide, Sn0 2 , as has been stated, is the principal form in which tin occurs in nature. The mineral is known as cassiterite or tin-stone. It is found at Cornwall, England; on the East Indian islands Banca and Biliton; and in Malacca. It crystallizes in the tetragonal system, and is generally colored from brown to lack. It is formed as a white powder by burning tin in the air, and by heat- ing the different varieties of stannic hydroxide. It is infusible, and is not acted upon by concentrated hydro- chloric or nitric acid. Concentrated sulphuric acid, how- ever, converts it into a gelatinous liquid from which stannic oxide is precipitated by water. Stannous Sulphide, SnS, which is formed when hy- drogen sulphide is passed into a solution of stannous chloride, is a brownish-black powder. When treated with the soluble sulphides and sulphur, or with a soluble poly- sulphide, it dissolves, forming a sulphostannate as repre- sented in these equations : SnS + (NH 4 ) 3 S + S = (NH 4 ) 2 SnS 3 ; SnS + (NHJA = (NH 4 ) 2 SnS 3 . Stannic Sulphide, SnS 2 .— This compound is obtained in crystallized form by heating together tin-filings, sulphur, 5^2 COLLEGE CHEMISTRY. and dry ammonium chloride; and in amorphous form by treating a solution of stannic chloride with hydrogen sul- phide. In the former case it is a golden-yellow crystalline substance; in the latter a yellow powder. The crystalline variety is known as mosaic gold. When heated to a high temperature it is converted into stannous sulphide. The precipitated variety is easily dissolved by concentrated hydrochloric acid, and converted into metastannic acid by concentrated nitric acid. The crystallized variety is not soluble in hydrochloric acid, and is affected but slightly by nitric acid. The crystallized variety, or mosaic gold, is used as a pigment, particularly for bronzing. Stannic sulphide dissolves easily in the soluble sulphides, forming sulphostannates : SnS, + (NH 4 ) 2 S = (NH 4 ) 2 SnS 3 ; SnS, + K 2 S = K s SnS 3 . Reactions which are of Special Value in Chemical Analysis. — The reactions with hydrogen sulphide, and the conduct of the precipitated sulphides when treated with ammonium sulphide or polysulphide, are constantly utilized for analytical purposes when tin is present. The reducing action of stannous compounds serves to distinguish them from stannic compounds. The carbonates of the alkali metals precipitate the hydroxides, and these do not dissolve in an excess of the carbonate. Metallic zinc precipitates the tin from a solu- tion of a tin compound as a spongy mass. If the precipi- tation is carried on in a platinum vessel, the platinum is not colored by it. By this means it can be distinguished from antimony, which is reduced under the same circum- stances, but is deposited as a black coating upon the platinum. Lead, Pb (At. Wt. 206.9). General. — The basic properties of lead are stronger, and its acidic properties weaker, than those of tin. Its prin- LEAD. 583 cipal compounds are those in which it acts as a base-form- ing element. The compounds in which it is quadrivalent, such as Pb0 2 , are comparatively unstable, and when treated with acids they readily pass over into the compounds of the series in which the lead is bivalent. Thus lead oxide itself readily gives up oxygen when treated with acids, and yields salts of bivalent lead. Forms in which Lead occurs in Nature. — Lead occurs in nature as the sulphide, PbS, which is known as galenite or galena. Other natural compounds of the metal are the carbonate, PbC0 3 , known as cerussite; the phosphate, Pb 3 (P0 4 ) 2 ; the chromate, Pb0rO 4 , or crocoisite; and the molybdate, PbMo0 4 , or wulfenite. Metallurgy. — Most of the lead in the market is obtained from the sulphide, and as most of the sulphide contains silver both metals have to be considered in the treatment of the ore. Under the head of Silver (which see) reference was made to the methods by which this metal is separated from the lead after both have been separated from their compounds. Here it will only be necessary to show how the metals are extracted together from the ore. This is accomplished in one of two ways : (1) By heating the sulphide with iron, when the latter combines with the sulphur, forming iron sulphide, while the lead is set free. (2) By roasting the sulphide until it is partly converte into lead oxide and lead sulphate, and then heating th mixture without access of air, when two reactions takt place, which are represented in these equations: PbS + 2PbO = 3Pb + S0 2 ; PbS + PbS0 4 = 2Pb + 2SO a . The lead is thus set free, and the sulphur is driven off as sulphur dioxide. Properties. — Lead is a bluish-gray metal, with a high lustre. It is very soft, and not very strong; melts at 325°, and has the specific gravity 11.37. At a high tem- perature it is converted into vapor. When heated in con- 5^4 COLLEGE CHEMISTRY. tact with the air it becomes covered with a layer of oxide as can be seen in a vessel containing the molten metal. Pure water acts upon lead when air has access to it, and some of the lead dissolves. If the water contains salts in solution, such as calcium carbonate, gypsum, etc. , or if it contains carbon dioxide, it acts only very slightly upon the metal; further, water that contains organic matter in a state of decomposition dissolves lead with comparative case. Nitric acid dissolves lead readily; as, however, lead nitrate is insoluble in nitric acid, it is necessary to use comparatively dilute acid. Concentrated sulphuric acid dissolves lead to some extent, and therefore a little lead sulphate is sometimes contained in commercial sulphuric acid. "When such a solution is diluted with water the sulphate is precipitated. Hence when some specimens of the commercial acid are diluted with water they become turbid, and after standing for a time they become clear, as the lead sulphate settles. Hydrochloric acid acts only slightly upon lead. Acetic acid dissolves the metal very readily. It is precipitated in metallic form from a solu- tion of one of its salts by metallic zinc. The formation is sometimes very beautiful, especially if the zinc is sus- pended in the solution. It is called the " lead tree," or Arbor Satumi. The action consists in a substitution of zinc for the lead. After the action is complete all the lead is deposited as metallic lead, and the zinc has entered into its place, forming a salt which remains in solution. Thus, if lead nitrate is used, zinc nitrate is in the solution. Applications. — Lead is extensively used for a variety of purposes, as, for example, for making sulphuric-acid chambers, for evaporating-pans for alum and sulphuric acid, for shot, for water-pipes, and in making various alloys. The use of lead water-pipes is a matter of much importance from the sanitary point of view, as is evident from the statements above made concerning the action of water upon the metal. Ordinary drinking-water acts under most circumstances only very slightly upon lead, and not enough is dissolved to be dangerous to those using LEAD CHLORIDE, ETC. 5 8 5 the water. At the same time circumstances may at any time arise which will increase the solvent power of the water, and thus cause serious results; and it would un- doubtedly be better if the use of such pipes could be entirely avoided in cases in which the water is to be used for drinking purposes. Babbitt's metal, which is used for machine bearings, contains copper, antimony, ,tin, and lead. Lead Chloride, PbCl 2 , is formed when hydrochloric acid or a soluble chloride is added to a cold solution of a lead salt, and appears as a white precipitate. It is soluble in hot water, and is deposited in the form of long, needle- shaped crystals when the solution cools. It occurs in nature in small quantity, as cotunnite. Lead Tetrachloride, PbCl 4 , is a liquid that forms a crystalline mass at — 15°. It is decomposed into the dichloride and chlorine at the ordinary temperature. At 105° this takes place with explosion. It is formed by passing chlorine into concentrated hydrochloric acid con- taining lead chloride. After saturation, ammonium chlo- ride is added, when the double salt, PbCl 4 . 2NH 4 C1, separates in yellow crsytals. Ice-cold concentrated sul- phuric acid decomposes it into ammonium sulphate, hydrochloric acid, and lead tetrachloride, which collects as a yellow oil below the sulphuric acid. Lead Iodide, Pbl 2 , is a yellow substance that crystallizes from water in beautiful lustrous laminse. It is precipi- tated when potassium iodide is added to a solution of a lead salt. Oxides of Lead.— Lead forms five distinct compounds with oxygen, the formulas and names of which are as fol- lows: had suboxide, Pb 2 0; lead oxide, PbO; lead ses- quioxide, Pb 2 O s ; minium, Pb 3 4 ; and lead peroxide, Pb0 2 . Lead Suboxide, Pb 2 0. — This compound is formed when lead is exposed to the air. Lead Oxide, PbO.— This is formed by heating lead nitrate, and is then left behind in the form of a yellow powder. If heated to fusion it solidifies, forming a 5^6 COLLEGE CHEMISTRY. yellowish or reddish mass known as litharge. This is formed in large quantity in the process of separating silver from lead. With the strongest bases it forms salts similar to those formed by zinc. This is seen in the solubility of the hydroxide in sodium and potassium hydroxide, which is due to the formation of compounds known as plumbites : Tb(OH) 2 + 2KOH = Pb(OK) 2 + 2H a O. Heated with silicon dioxide it forms a silicate which is easily fusible. Lead oxide is used extensively in the manufacture of flint glass for optical purposes, as was described under Glass (which see). It also finds applica- tion in glass painting and porcelain painting, and is used in the manufacture of lead salts, particularly "sugar of lead," which is an acetate. Lead Peroxide, Pb0 2 , is formed by treating minium or red lead with dilute nitric acid. Minium has the compo- sition Pb s 4 . When treated with nitric acid, a part dis- solves as lead nitrate, and lead peroxide remains behind, as represented in the equation Pb s 4 + 4HN0 3 = Pb0 2 + 2Pb(N0 3 ) 2 + 2H 2 0. The peroxide is formed in general by the action of oxidiz- ing agents upon the lower oxides of lead. One of the most convenient methods of making it consists in treating ead acetate with a filtered solution of bleaching-powder. i is a dark-brown powder, insoluble in water. When ignited it loses half of its oxygen, and it gives up its oxygen readily to other substances. Towards hydrochloric acid it acts like manganese dioxide, giving lead chloride and chlorine according to the equation Pb0 2 + 4H01 = Pb01 2 + 2H 2 + Cl 2 . It appears probable that the tetrachloride is first formed, and that this then breaks down into the dichloride and chloride. When the peroxide is treated in the cold with hydrochloric acid it dissolves, and when this solution is STORAGE BATTERY. 587 heated it gives off chlorine. Further, when it is treated with caustic alkalies lead peroxide is thrown down. Lead peroxide dissolves in concentrated caustic potash and forms a salt of the formula K a Pb0 3 , or PbO<9S, analogous to potassium stannate, K 2 SnO s , silicate, K,SiO s , carbonate, K 2 CO s , etc. Other salts derived from the acid PbO(OH), are known, and are called plumbates. Storage Battery — An efficient galvanic battery can be made by putting two lead plates in dilute sulphuric acid, one of the plates being covered with a layer of lead peroxide. The sulphuric acid acts upon the lead, convert- ing it into lead sulphate, which forms a solid layer upon the plate. On the other hand, the acid acts upon the 'lead peroxide, reducing the lead to the bivalent condition and forming with it lead sulphate. As long as the chem- ical action can take place, that is to say, as long as there is lead peroxide present, the battery continues to act. When it ceases to act it can be restored to the active condition by passing an electric current through it in the opposite direction. Under these conditions the lead sulphate on one plate is reduced to metallic lead, and that on the other is oxidized to lead peroxide. These facts are taken advantage of in the preparation of the well-known storage batteries. Red Lead, Minium. — When lead oxide is gently heated in the air it takes up oxygen, and is converted into the red compound known as minium or red lead. The commer- cial article of this name varies in composition, but approximates to that represented by the formula Pb s 4 , and, if the oxide is slowly heated, the amount of oxygen taken up is that required to form a compound of the above formula. Red lead varies in color from red to yellowish, according to the method of preparation. When heated it becomes dark, but the red color appears again on cool- ing. When heated to a high temperature it loses oxygen and yields lead oxide: Pb 3 4 = 3PbO + 0. 588 COLLEGE CHEMISTRY. When treated with dilute nitric acid lead nitrate is formed, and lead peroxide is left undissolved. Red lead is used as a pigment, and sometimes in place of litharge when an oxide of lead is needed, as in the manufacture of glass, as a flux in the manufacture of porcelain, etc. Lead Sulphide, PbS. — This has already been referred to as the principal compound from which lead is obtained. The natural variety is called galena or galenite. It is formed in the laboratory as a black precipitate, when hydrogen sulphide is passed into a solution of a lead salt. When heated in the air, as in the roasting of galenite, the sulphur passes off as sulphur dioxide, and the lead is con- verted into oxide. Concentrated hydrochloric acid dis- solves it. Concentrated nitric acid converts it into the sulphate. Lead Nitrate, Pb(NO s ) 2 .— The nitrate is easily made by dissolving lead, lead oxide, or carbonate in nitric acid. The salt crystallizes well, and is easily soluble in water. It is difficultly soluble in dilute nitric acid, and insoluble in concentrated nitric acid, resembling in this respect barium nitrate. It is decomposed by heat, giving nitrogen peroxide, NO, , and lead oxide. Lead Carbonate, PbCO s . — The carbonate occurs in nature as cerussite, crystallized in forms which are the same as those of barium carbonate, and of that variety of calcium carbonate known as aragonite. It can be obtained by adding lead nitrate to a solution of ammonium car- bonate, but, when solutions of lead salts are treated with the secondary carbonates of the alkali metals, precipitates of basic carbonates are always obtained. When an excess of sodium carbonate is added to a solution of lead nitrate, the precipitate has the composition 3Pb0.2C0 2 + H,0. The salts usually obtained are more complicated than this, but the relations between them and lead oxide and car- bonic acid are of the same kind. Basic lead carbonate is prepared and used extensively, under the name of white lead, as a pigment. It is manufactured by different methods. The principal ones are the following : LEAD CARBONATE- LEAD SULPHATE. 589 (1) Hie Dutch Method. — This consists in exposing sheets of lead wound in spirals to the action of vinegar, air, and carbon dioxide from decaying organic matter. The spirals of sheet lead are placed in earthenware vessels, on the bottom of which, but not in contact with the lead, the vinegar is placed. The vessels thus arranged are placed in beds of horse-manure. In consequence of decomposi- tion, which is set up in the manure, carbon dioxide is given off slowly, and enough heat is generated to start the action upon the lead. The chemical changes involved in the process are, mainly, the formation of a basic acetate of lead, and the subsequent decomposition of this by carbon dioxide, forming a basic carbonate, and leaving the acetic acid free to act upon a further quantity of lead. _. (2) The French Method. — In this method a solution of basic lead acetate is prepared by treating a solution of the neutral salt with lead oxide. This is then decomposed by passing carbon dioxide into it, when a basic carbonate is thrown down. The carbon dioxide is generally made by burning coke. (3) The English Method. — This is a modification of- the Dutch method, and differs from it chiefly in the substitu- tion of spent tan in a state of fermentation for manure, and the use of dilute acetic acid in place of vinegar. There is less risk of discoloration in consequence of the formation of hydrogen sulphide, but the fermentation takes place more slowly, and the whole process, therefore, requires a longer time. ' ' An objection to white-lead paint is thit it turns dark under the influence of hydrogen sulphide. It also turns yellow in consequence of the action of some substance contained in the oil with which the lead carbonate is mixed. Lead Sulphate,, PbS0 4 , occurs to some extent in nature. It is formed by adding sulphuric acid or a soluble sulphate to a solution of a lead salt, and by oxidation of lead sul- phide. Like barium sulphate, it is practically insoluble in water. As stated above, it is somewhat soluble in con- 59° COLLEGE CHEMISTRY. centrated sulphuric acid, and it is therefore always found in the concentrated acid of commerce. Nitric acid and hydrochloric acid dissolve it in considerable quantity. It dissolves further quite readily in solutions of some am- monium salts, as in ammonium tartrate and acetate. When heated to redness it is partly decomposed with loss of sulphur trioxide. Reactions which are of Special Value in Chemical Analysis — .The reactions of lead salts with the soluble hydroxides, with sulphuric acid, hydrochloric acid, hy- drogen sulphide, soluble carbonates, potassium chromate and dichromate, are the ones which are principally used in analysis. All of these have been treated of in this chapter, with the exception of those with potassium chromate and bichromate, which will be taken up in the chapter on Chromium (which see). In anticipation it may be said that the reactions are based upon the fact that lead chromate, PbCr0 4 , like barium chromate, is in- soluble in water. The elements of Family V, Group A, are vanadium, columbium, didymium, and tantalum. As they are closely related to the members of Group B, of the same family, they were treated of at the end of Chapter XVIII in con- nection with the members of the phosphorus group. Among them the one which is least known is didymium. This in turn is more or less closely related to two other elements of nearly the same atomic weight which occur in Families III and IV. These are lanthanum and cerium. A few words in regard to these three rare elements will suffice for the present purpose. Lanthanum, Cerium, Didymium. These three elements occur together in several rare minerals of Norway, as cerite, gadolinite, and allanite. Cerite is a silicate of the three metals, and its composition LANTHANUM. ETC. 59 1 * i is represented by the formula La 4 \ (SiO,)_ -f- 3H 0. It DiJ is probably a mixture of three isomorphous silicates. The principal constituent is cerium silicate, Ce 4 (SiOJ s . The perfect separation of the constituents of the mineral is a very difficult operation. Lanthanum, La (At. Wt. 138), forms an oxide of the formula La a O s , analogous to that of aluminium. Its chloride also is analogous to that of aluminium, and has the composition LaCl,; and in all its salts it acts as a trivalent element. Cerium, Ce (At. Wt. 140), forms two series of com- pounds, in one of which it is trivalent, resembling lan- thanum, and the other members of the aluminium group; while in the other series it is quadrivalent, resembling silicon and the other members of the silicon group. Didymium, Di, has already been referred to on p. 339 in connection with the members of Family V, Group A, which it resembles in some respects. In most of its com- pounds it is, however, trivalent, forming compounds, of some of which the following are the formulas : DiOl,, Di a O„ Di(NO s ) s , Di^SO,),, Di a (C0 8 )„ etc. Praseodymium, Pr, and Neodymium, Nd. — While the name didymium is still given above, and this substance dealt with as though it were an element, as it was at first held to be, it has been shown by Auer von Welsbach that it consists of two very similar elements to which he has given the names praseodymium and neodymium. When the double nitrate of ammonium and didymium is re- peatedly recrystallized it is separated into two salts, one of which is green, and the other rose-colored. When the nitrate or oxalate of one of these new elements is ignited it forms a black oxide, while from the other is formed an oxide of a different color. The element that gives green salts is called praseodymium, and the other is called S9 2 COLLEGE CHEMISTRY. neodymium. The atomic weight of neodymium is 143.6; that of praseodymium 140. 5. EXPERIMENTS. Tin and its Compounds. -r Experiment 227. — Dissolve tin in hydrochloric acid and let the product crystallize. Experiment 228. — Pass dry chlorine over granulated tin con- tained in a retort connected with a receiver, using the arrange- ment illustrated in Fig. 60, page 318. Redistil the product. Treat some of the liquid with water, and boil. Lead and its Compounds. Experiment 229. -Make specimens of lead chloride and lead iodide, and crystallize them from water. Experiment 230. — Make lead sesquioxide by bringing together lead acetate and sodium hydroxide, and treating the solution with a solution of sodium hypochlorite. Experiment 231. — Treat some red lead with dilute nitric acid. Filter, wash, and treat the substance left on the filter with hydrochloric acid- CHAPTER XXXI. ELEMENTS OF FAMILY VI, GROUP A: CHROMIUM.— MOLYBDENUM.— TUNGSTEN- URANIUM. General — At the end of Chapter XIV, in connection with the elements of the sulphur group, the four elements which form the subject of this chapter were briefly referred to, for the reason that in some respects they resemble sulphur. As was there stated, this resemblance "is seen mainly in the formation of acids of the formulas H CrO , H 2 Mo0 4 , H,W0 4 , and H a U0 4 ; and the oxides "CrO*, MoO,, WO s , and U0 3 ." Further, it was stated that "when the acids of chromium, molybdenum, tungsten, and uranium lose oxygen, they form compounds that have little or no acid character. The lower oxides of chromium form salts with acids, and these bear a general resemblance to the salts of aluminium, iron, and manganese. The chromates lose their oxygen quite readily when acids are present with which the chromium can enter into combina- tion as a base-forming element." "Molybdenum and tungsten do not form salts of this character: indeed they seem to be practically devoid of the power to form bases. Uranium, on the other hand, forms some curious salts which differ from the simple metallic salts which we com- monly have to deal with. These are the uranyl salts which are regarded as acids, in which the hydrogen is either wholly or partly replaced by the complex U0 2 , which is bivalent. Thus the nitrate has the formula UO^NO,).,, the sulphate (U0 2 )S0 4 , etc. These salts are derived from the compound TJ0 2 (OH) 2 , acting as a base, whereas the compound has also distinctly acid properties." That 593 594 COLLEGE CHEMISTRY. member of the group the compounds of which are most commonly met with in the laboratory and in the arts is chromium, and this will receive principal attention here. • Chromium, Cr (At. "Wt. 52.1). General. — This element forms three series of com- pounds, in which it appears to be respectively bivalent, trivalent, and sexivalent. Of these the members of the series in which it is trivalent are most stable under ordinary circumstances. Some of the principal members of the first series, or the chromous compounds, are represented by the formulas CrCl 2 , Cr(OH) 2 , CrS0 4 , CrC0 3 . Of the second series, or the chromic compounds, some of the principal members are : CrCl„ Cr 2 3 , Cr 2 (S0 4 )„ Cr(NO,) s , KCr(S0 4 ), + 12H 2 0. And, finally, the members of the third series are derived from the oxide CrO, , and they are for the most part salts of the acid of the formula H 2 Cr0 4 , known as chromic acid, or of an acid of the formula H 2 0r 2 O, , known as bichromic acid, which is closely related to chromic acid. When exposed to the air the chromous compounds are converted into chromic compounds, and they are in gen- eral readily converted into chromic compounds by the action of oxidizing agents, as cuprous and mercurous compounds are converted into cupric and mercuric com- pounds. If the oxidation takes place in acid solution the limit is reached when a chromic salt is formed. If, how- ever, the action takes place in the presence of a strong base the limit is reached in the formation of a chromate. Thus, suppose chromous oxide to be treated with an oxidizing agent in the presence of sulphuric acid, the final CHROMIUM. 595 product would be chromic sulphate, as represented in the following equations: CrO + H a SO, = CrS0 4 + H a O; 2CrS0 4 + H a SO, + = Or,(S(\), + H a O. On the other hand, if the oxidation takes place in the presence of caustic potash the final product is potassium chromate, as shown in the following equation: CrO + 2KOH + 0, = K a Cr0 4 + H 2 0. When a chromate is treated with an acid it tends to pass back to a compound of the chromic series, and the change involves the giving up of oxygen. Forms in which Chromium occurs in Nature. — The principal form in which chromium occurs in nature is the mineral chromite, also known as chromic iron and chrome iron ore. This has the composition FeCr 2 4 , and, as will be pointed out below, it is probably analogous to the spinels (see p. 526), being an iron salt of the acid CrO. OH, which may be called metachromous acid. This view is represented by the formula «q - Q>Fe. It occurs also in the mineral crocoisite, which is lead chromate, PbCr0 4 . The name chromium is derived from the Greek xpcS/ta, meaning color; and the element is so called because most of its compounds are colored. Preparation.— Chromium is prepared by heating chromic oxide with carbon in the electric furnace. The first product obtained is a mixture of carbides of chromium. When this is heated with lime in the electric furnace calcium carbide and metallic chromium are formed. By this method it is obtained in crystals. Properties. — Chromium has a bright metallic lustre and is capable of high polish. It is very hard and difficultly fusible. It burns brilliantly in oxygen, though it is not changed by exposure to the air at ordinary temperatures. It is dissolved by hydrochloric acid. Nitric acid does not 596 COLLEGE CHEMISTRY. attack it. When treated with salts of potassium which easily give up their oxygen, as the chlorate and nitrate, it is converted into potassium chromate. Chromic Chloride, CrCl 3 . — This compound is made in solution by dissolving chromic hydoxide, Cr(OH) s , in hydrochloric acid. This solution has a dark-green color. When evaporated to a sufficient extent crystals of the com- position Cr01 3 -|- 6H 2 are deposited. If these are heated in the air they undergo decomposition just as aluminium chloride does, and the product left behind is chromic oxide : 2CrCl 3 + 3H 2 = Cr 2 3 + 6HC1. If, however, the crystallized chloride is heated in an atmosphere of chlorine or hydrochloric acid, the water is given off, and the anhydrous chloride, which has a beauti- ful reddish-violet color, is formed. This dissolves in water and forms a green solution. But if the dry chloride thus obtained is sublimed, it is deposited in lustrous laminae of the same color; and this variety is insoluble in water and acids, and is only slowly acted upon by boiling alkalies. This insoluble, crystallized variety of the chloride is obtained also by the same method as that used in making aluminium chloride, that is, by passing a current of chlo- rine over a heated mixture of carbon and chromic oxide. Chromic Hydroxide, Cr(OH) 3 . — When ammonia is added to a solution of a chromic salt, a light-blue volu- minous precipitate, which has the composition Cr(OH) 3 -|- 2H 2 0, is formed. When this is filtered off and dried in a vacuum it loses the water and leaves the hydroxide. This is readily converted by heat into a compound of the formula CrO. OH, and finally into chromic oxide, Cr 2 O s . The green precipitates formed in solutions of chromic salts by sodium and potassium hydroxides always contain some of the alkali metal in combination. Chromic hydroxide, like aluminium hydroxide, dissolves in the soluble hydroxides, and forms salts known as chromites. which are derived from the acid CrO. OH. CHROMIC OXIDE— CHROMIC SULPHATE. 597 Thus with potassium hydroxide the action takes place as represented in the equation Cr^-OH + KOH = C< jj K + 2H 2 0. When a solution containing potassium or sodium chromite is boiled the salt is decomposed and chromic hydroxide precipitated, though the precipitate thus formed always contains some of the alkali metal in combination. It will be noticed that in this respect aluminium and chromium conduct themselves differently towards the alkaline hy- droxides. It has already been stated that chromite, (CrO.O) 2 Fe, is regarded as an iron salt of the same order as the potas- sium salt referred to. Chromic Oxide, Cr 2 3 , is formed by igniting the hy- droxides, and is most readily prepared by heating a mixture of potassium dichromate and sulphur. The sulphur is oxidized, and with the potassium forms potassium sul- phate, while the chromic acid is reduced to the form of the oxide Cr 2 3 . It can be obtained in crystals. As ordinarily obtained it is a green powder, which after igni- tion is almost insoluble in acids. It is, however, dissolved by treatment with fusing acid potassium sulphate. The oxide colors glass green, and is used in painting porcelain. Chromic Sulphate, Cr 2 (S0 4 ) 3 , is made by dissolving the hydroxide in concentrated sulphuric acid when it is deposited in purple crystals of the composition Cr 2 (SO t ) 3 -+- 15H 2 0. If the solution of this salt is boiled, the solution becomes green, and crystals cannot be obtained from it. But by standing for some time the green solution becomes reddish purple again, and yields the crystallized salt. Other salts of chromium act in the same way. They exist in two varieties, one of which crystallizes and is reddish- purple in color, while the other does not crystallize and is green. The crystallized salts are converted into the un- crystallized green salts by boiling, and the green salts are converted into the crystallized salts by standing. 598 COLLEGE CHEMISTRY. Chrome-Alums.- — Chromic sulphate, like aluminium sulphate, combines with other sulphates, such as potas- sium, sodium, and ammonium sulphates, and forms well- crystallized salts, which are closely, analogous to ordinary alum. They all contain twelve molecules of water, as represented in the formulas below: r Chrome- Alum KCr(S0 4 ) 2 + 12H,0 Sodium Chrome- Alum NaCr(S0 4 ) 2 + 12H 2 Ammonium Chrome-Alum (NHJCr(S0 4 ) 2 + 12H 2 The potassium compound which is commonly called chrome-alum is made by adding a reducing agent, such as alcohol or sulphur dioxide, to a solution of potassium dichromate containing sulphuric acid. If the solution is heated it turns green, and crystals cannot be obtained from it. But on standing for a considerable time its color changes, and reddish-purple crystals of the alum are deposited. This change can be facilitated by putting some crystals of the salt in the concentrated green solu- tion. The action of reducing agents upon potasbium dichromate will be treated of farther on. The salt finds application in dyeing and tanning. " Chromic Acid and the Chromates — It has been stated / that when chromium compounds belonging to the chro- U mous and chromic series are oxidized in the presence of V bases they are converted into chromates. These salts are derived from an acid of the formula H 2 Cr0 4 , which is unknown, as it breaks down spontaneously into chromium trioxide, Cr0 3 , and water, when it is set free from its salts, just as carbonic and sulphurous acids break down respectively into carbon dioxide and water, and sulphur dioxide and water. The starting-point for the preparation of the chromates and the compounds related to them is chromic iron. This is ground fine, intimately mixed with a mixture of caustic potash and lime, and then heated in shallow furnaces in contact with the air. Under these circumstances oxidation is effected by the oxygen of the air. The iron is converted into ferric oxide, and the / / POTASSIUM CHROMATE, ETC. 599 chromium gives, with the calcium and potassium, the corresponding chromates, CaCr0 4 and K 2 0rO 4 . When the mass is treated with water these salts dissolve, and ferric oxide remains undissolved. By treating the solution with potassium sulphate the calcium salt is converted into the potassium salt, and thus all the chromium appears in the form of potassium chromate. The changes referred to are represented in the following equations : / 2(Cr0 2 ) 2 Fe + 8KOH + 70 = 4K 2 Cr0 4 -f Fe 2 3 + 4H 2 0; 2(0r0 1 ) 1 Pe + 4CaO + 70 = 4CaCr0 4 + Fe 2 O f ; CaCr0 4 -t-K 2 S0 4 = K 2 Cr0 4 + CaS0 4 . \ As potassium chromate is easily soluble in water, and therefore difficult to purify, it is converted into the bichro- mate, which is less soluble and crystallizes well. The change is easily effected by adding the necessary quantity of a dilute acid. If nitric acid is used the reaction is represented by the following equation : 2K 2 Cr0 4 + 2HN0 3 = K 2 Cr 2 0, + 2KN0 3 + H 2 0. The salt thus obtained is manufactured on the large scale and is the starting-point for the preparation of other chromium compounds. Potassium Chromate, K 2 Cr0 4 , formed as above described, is a light-yellow crystallized substance which is easily soluble in water. It is isomorphous with potassium sulphate. Acids convert it into the bichromate, as just stated. Potassium Bichromate, K 2 Cr 2 0,. — This salt forms large triclinic red plates. It is soluble in ten parts of water at the ordinary temperature, and is much more soluble in hot water. When heated it at first melts without undergoing decomposition; at white heat, however, it is decomposed, yielding the chromate, chromic oxide, and oxygen : 2K,Cr 2 7 = 2K 2 Cr0 4 + Cr 2 3 + 30. 600 COLLEGE CHEMISTRY, It undergoes a similar change, but much more readily, when heated with concentrated sulphuric acid. In this case, however, the chromic oxide forms chromic sulphate with the acid, and this forms chrome-alum with the potas- sium sulphate: K 2 Cr 2 7 + 4H,S0 4 = 2KCr(SOJ 2 + 4H a O + 30. All the oxygen in the chromate in excess of that required to form the alum and water is given off. This also is the character of the action towards reducing agents in general. Concentrated hydrochloric acid is oxidized by the dichromate, and chlorine is evolved : K 2 Cr 2 7 + 14HC1 = 2KC1 + 2CrCl 3 + 7H a + 601. Here two atoms of chlorine are required to form potassium chloride with the potassium, and six to form chromic chloride with the chromium; and the eight hydrogen atoms in combination with this chlorine combine with four atoms of oxygen of the dichromate, leaving three more to oxidize hydrochloric .acid. Consequently one molecule of the dichromate sets free six atoms of chlorine : 30 + 6HC1 = 3H 2 + 6C1. When the bichromate in solution is treated with potas- sium hydroxide, its color changes to yellow, in consequence of the formation of the chromate, the action taking place as represented in this equation : K 2 Cr a O, + 2K0H = 2K 2 CrQ 4 + H a 0. Potassium bichromate finds extensive use in the arts and in the laboratory as an oxidizing agent. With gelatin it forms a mixture sensitive to light, which turns it dark, and makes it insoluble. This fact is made the basis of a number of photogravure or photographic reproduction processes. The bichromate is used, further, in dyeing. Chromium Trioxide, CrO s . crystallizes out on cooling when either the chromate or the bichromate is treated in CHROMATES AND BICHROMATES. 60 1 concentrated solution with concentrated sulphuric acid. It is a beautiful red substance, which crystallizes in needles. When dissolved in water it. forms a solution from which, by neutralization, the chromates can be obtained. It is an extremely active oxidizing agent, disintegrating most organic substances with which it is brought in contact. Relations between the Chromates and Bichromates. — The fact that chromium trioxide with water gives chromic acid, which is a dibasic acid, whose salts in general resem- ble those of sulphuric acid, leads to the belief that the structure of chromic acid should be represented by a formula similar to that of sulphuric acid, thus: O s S(OH), 2 Cr(OH) 3 . Just as sulphuric acid by loss of water is converted into disulphuric acid or pyrosulphuric acid, so chromic acid is converted into bichromic acid, and in all probability the relation between the chromates and bichromates is the same as that between the sulphates and disulphates, as represented by the equations SO,g" SO >.<0H But, as has been stated, neither chromic acid nor bi- chromic acid is known, as they break down into chromium trioxide and water when set free from their salts. The conversion of potassium chromate into the bichromate by treatment with an acid is represented as shown below: Cr0 *so * = uo . so > + 3H 3 °; TJO < 0H 4. H0 - N0 * - UO <°- m * 4- 2H O uu »Ca, is formed by heating together ferric oxide and lime to a high temperature. In composi- tion this is plainly analogous to the spinels. Magnetic oxide of iron or magnetite is believed to be the correspond- ing ferrous salt, ^ < -.' ( - ) >Fe. Franklinite also is a salt of the same order, containing zinc. It is essentially a zinc salt, of the formula -r, Q - Q>Zn, but some of the zinc is replaced by iron and manganese. Ferrous-Ferric Oxide, Fe 3 4 . — As stated above, this compound is regarded as analogous to the spinels, and as the ferrous salt of the acidic hydroxide FeO.OH, as repre- sented in the equation (FeO. 0) 2 Fe. It is found in nature as the mineral magnetite, and loadstone, which occurs in Sweden, Norway, and elsewhere. It is, further, formed when iron is burned in oxygen, and when water is passed over red-hot iron. Some of the magnetite which occurs in nature has the power to attract iron, or is magnetic. Soluble Ferric Hydroxide is formed when a solution of ferric chloride or ferric acetate is treated with ferric hydroxide, and the solution thus formed dialyzed (see p. 395). The ferric salts pass through the membrane, and the ferric hydroxide remains in solution in water, 632 . COLLEGE CHEMISTRY. forming a deep-red liquid. It is used in medicine. Small quantities of salts cause the precipitation of ferric hy- droxide from the solution. Ferric Oxide, Fe 2 3 , is found in nature, and is known as hematite, forming one of the most valuable ores of iron. It can be made in the laboratory by igniting the hydroxide. As hematite, it is a black, crystallized substance with a high lustre. Otherwise it has a red or a reddish-brown color. The oxide found in nature and that which has been strongly ignited are very difficultly soluble in acids. In the preparation of fuming sulphuric acid by heating ferrous sulphate (see p. 232) there is left a residue of ferric oxide known as rouge, which is used as a red pigment and as a polishing powder. A specially fine variety of rouge for polishing is manufactured by heating ferrous oxalate, FeC 2 4 , in contact with the air. Ferrous Sulphide, FeS, is formed by direct union of iron and sulphur when the two are heated together. It is manufactured by heating iron filings and sulphur together in a crucible. The pure compound is yellow and crystalline. When heated in contact with the air it is oxidized to ferrous sulphate, if the temperature is not too high. At a higher temperature the products are sulphur dioxide and ferric oxide. When a solution of a ferrous salt is treated with ammonium sulphide, ferrous sulphide is precipitated as a black powder. When a ferric salt is treated with ammonium sulphide it is reduced to the ferrous condition, and then ferrous sulphide is precipi- tated : Fe 2 (S0 4 ) 3 + (NH 4 ) 2 S = 2FeS0 4 + (NHJ.SO, -f S; 2FeS0 4 . + 2(NH 4 ) 2 S = 2FeS + 2(NH 4 ) 2 S0 4 . The sulphide thus obtained is readily oxidized in the air, and forms the sulphate. The compact variety is used in making hydrogen sulphide (which see). Ferrous Carbonate, FeC0 3 . — This salt occurs in nature as spathic iron or siderite. It crystallizes in forms similar FERROUS SULPHATE, ETC. 633 to those of calc spar or calcium carbonate, CaC0 3 . Like this, further, it dissolves in water which contains carbon dioxide, and is therefore contained in natural waters which come in contact with it. When a solution of a ferrous salt is treated with a soluble carbonate a white precipitate is formed, which is ferrous carbonate; but in contact with the air this is rapidly oxidized and decomposed, leaving ferric hydroxide, which with carbonic acid does not form a salt. In this respect ferric hydroxide acts like aluminic and chromic hydroxides, and therefore when a soluble carbonate is added to a solution of a ferric salt the hydroxide and not ferric carbonate is thrown down. Ferrous Sulphate, FeS0 4 . — This important compound is manufactured on the large scale by the spontaneous oxidation of pyrite in contact with the air, and by dissolv- ing iron in sulphuric acid. It is frequently called "green vitriol" (see p. 544), and more commonly "copperas." Under ordinary conditions it crystallizes in transparent, green, monoclinic crystals with seven molecules of water, just as zinc sulphate, magnesium sulphate, etc., do; and when heated, six of these are given off readily, while the last is given off with difficulty — a fact which makes it appear probable that the salt is a derivative of tetra- hydroxyl-sulphuric acid, as represented in the formula ( (OH), OS < -n . While it ordinarily crystallizes in monoclinic crystals, it takes the rhombic form if its supersaturated solution is touched with a crystal of zinc sulphate. It also crystallizes in the triclinic system with five molecules of water, like cupric sulphate, if a crystal of the latter salt is placed in its concentrated solution. Ferric Sulphate, Fe 2 (SOJ 3 . — This salt is formed by oxidation of ferrous sulphate. It is also formed by dis- solving ferric oxide or hydroxide in sulphuric acid. Ferrous Phosphate, Fe 3 (P0 4 ) 2 , occurs in nature crystal- lized with eight molecules of water as the mineral vivianite. Both this salt and ferric phosphate, FeP0 4 , are insoluble 6 34 COLLEGE CHEMISTRY. and are formed when solutions of ferrous and ferric salts are treated with sodium phosphate. Iron Disulphide, FeS 2 , is not analogous to any oxygen compound of iron. In it the metal appears to be quad- rivalent. The disulphide occurs very widely distributed and in large quantities in nature as the mineral iron pyrites or pyrite, which crystallizes in the regular system, and as marcasite, which crystallizes in the rhombic system. It can be made artificially, and if crystallized it appears in the form of pyrite. Its conduct under the influence of heat has been repeatedly referred to in connection with the roasting of iron and other ores. As pyrite it has a golden-yellow color, and it has frequently been taken for gold by those not familiar with it. The name "fool's gold," by which it is sometimes popularly known, suggests this fact. Arsenopyrite, FeAsS, also called arsenical pyrites, occurs in nature, and is a valuable source of the element arsenic; for, as has been stated (see p. 305), when it is heated it gives off arsenic, and ferrous sulphide is left behind. Reactions which are of Special Value in Chemical Analysis. — Ferrous Compounds. — The reactions of ferrous compounds with the soluble hydroxides and carbonates, ammonium sulphide, potassium ferricyanide, and with oxidizing agents have been explained above. With am- monium salts ferrous chloride forms double salts, which are soluble; therefore, if ammonium chloride is added to a solution of the salt ammonia does not precipitate the hydroxide. Further, ammonia does not completely pre- cipitate the hydroxide from a solution of a ferrous salt, as an ammonium salt is formed. By standing in the air, however, these solutions containing the double salts are oxidized, and ferric hydroxide is precipitated. The reac- tions with potassium cyanide will be understood from what" has been said concerning the compounds of ferrohydro- cyanic and ferrihydrocyanic acids. Ferric Compounds. — The reactions of ferric compounds with the soluble hydroxides and carbonates, ammonium COB A IT. 635 sulphide, potassium ferrocyanide, and potassium ferri- cyanide have been explained above. When hydrogert\ sulphide is passed through a solution of a ferric salt, \ reduction to the corresponding ferrous salt takes place, / and sulphur separates, which gives the solution a milky"^ appearance : Fe 2 (S0 4 ) 3 + H 2 S = SFeSO, + H 2 S0 4 + S; 2FeCl 3 + H 2 S = 2FeCl 2 + 2HC1 + S. When a neutral solution of a ferric salt is treated with suspended barium carbonate the iron is precipitated as the hydroxide. When a neutral solution of a ferric salt is treated with acetate of potassium or sodium it turns dark red, in conse- quence of the formation of ferric acetate which remains in solution. When the solution is boiled the acetate breaks down into acetic acid and a basic hydroxide of iron, which is precipitated. When the basic acetate is filtered off, washed, and heated it is changed to ferric hydroxide: FeCl 3 + 3NaC 2 H 3 2 = Fe(C 2 H 3 2 ) 3 + 3NaCl; Fe(C 2 H 3 2 ) 3 + 3H.0 = Fe(OH) 3 + 3C 2 H 4 2 . When potassium sulphocyanate, KCNS, is added to a solu- tion of a ferric salt a blood-red color is produced. This occurs even in extremely dilute solutions of ferric salts. The borax bead is colored bottle-green in the reducing flame, and brown-red to yellowish red in the -oxidizing flame, when treated with compounds of iron. Cobalt, Co (At. Wt. 59). General. — As stated in the remarks introductory to the iron group, cobalt, like nickel, preferably forms com- pounds which are analogous to ferrous compounds. It, however, forms a few which are analogous to ferric com- pounds, its power in this direction being greater than that 636 COLLEGE CHEMISTRY. of nickel. Its salts form a great variety of compounds with ammonia, and these have been extensively studied. Occurrence and Preparation — Cobalt occurs in nature, almost always in company with nickel. The principal minerals containing it are smaltite, CoAs 8 , and cobaltite, CoS 2 .CoAs 2 . In each of these iron, and generally some nickel, take the place of a part of the cobalt. Properties. — Cobalt has a silver-white color, with a slight cast of red. It is harder than iron, and melts at a somewhat lower temperature; is tenacious; and has the specific gravity 8.9. It dissolves in nitric acid. Cobaltous Chloride, CoCl 2 , is formed by heating cobalt in chlorine gas, and in solution by treating cobalt car- bonate with hydrochloric acid. From the solution it crystallizes in dark-red prisms of the composition CoCl 2 -f- 6H 2 0. The anhydrous salt is blue. When the blue salt is treated with water it turns red, and when the red salt is heated it turns blue. This difference in color between the anhydrous and the hydrated salts is charac- teristic of cobalt salts. If marks are made on paper with a dilute solution of one of the salts the color is not per- ceptible. If, however, the paper is held before a fire, the salt loses water and turns blue, and as the blue is more intense than the red, it is visible. When the salt becomes moist again it becomes invisible. This is the basis for the preparation of the so-called sympathetic inks. Cobaltous Hydroxide, Co(OH) 2 , is formed as a red pre- cipitate when a soluble hydroxide is added to a cobaltous salt, and the blue precipitate, which is first formed and which is a basic salt, is allowed to stand. It is oxidized by contact with the air, forming cobaltic hydroxide, which breaks down into cobaltic oxide, Co 2 3 . Cobaltic Hydroxide, Co(OH) 3 , is formed when calcium hypochlorite is added to a solution of a cobaltous salt, and is a black powder. When heated it is converted into black cobaltic oxide, Co 2 O s . Cobalt Sulphide, CoS, is the black precipitate which is formed by adding ammonium sulphide to a cobaltous salt. NICKEL. 637 It is not soluble in dilute acids, and differs from ferrous sulphide in this respect. Other sulphides of cobalt are those of the formulas Co 3 S 4 and CoS a . The former is found in nature, and is known as linnaeite. The latter occurs m combination with other sulphides, and with arsenides; as in cobaltite, CoS 2 .CoAs 2 . Cyanides. — Cobaltous cyanide is an insoluble dirty-red compound which is formed when potassium cyanide is added to a solution of a cobalt salt. It dissolves in an excess of potassium cyanide, forming a double cyanide, K 4 Co(ON) 6 , which is analogous to potassium ferrocyanide. When this solution is boiled the cyanide is oxidized, form- ing a compound analogous to potassium ferricyanide, thus : 2K 4 Co(CN) 6 + H 2 + = 2K 3 Co(CN) 6 + 2KOH. This acts like the corresponding iron compound. The cobalt is not precipitated from it by ammonium sulphide or sodium hydroxide. This conduct towards potassium cyanide distinguishes cobalt from nickel salts. Smalt. — The beautiful pigment known by this name is essentially a cobalt glass in which cobalt takes the place of calcium. It is made by heating compounds of cobalt with quartz and potassium carbonate. The glass thus formed is powdered very finely and used as a pigment. It does not change color in the sunlight, and is not affected by acids nor by alkalies. Chester J. Hunn Nickel, M (At. Wt. 58.7). General. — Nickel differs from cobalt in respect to the difficulty with which it forms nickelic compounds or those in which it is trivalent. At the same time it does form an oxide of the composition Ni 2 3 , and the corresponding hydroxide, Ni(OH) 3 . In all other compounds it is bivalent, the compounds being analogous to ferrous compounds. Occurrence and Preparation. — Nickel occurs native in meteorites. The principal minerals containing it are the 638 COLLEGE CHEMISTRY. arsenide, NiAs, known as niccolite, and the sulpharsenide, NiSAs or NiS 2 . NiAs a , known as gersdorffite. Properties. — Nickel is a white metal with a slight cast of yellow. It is very hard, and capable of a high polish. The metal in its ordinary condition is brittle, but after treatment with a little magnesium it becomes very mal- leable. Its specific gravity is 8.9, and it melts at a high temperature. It is not changed in the air; it dissolves slowly in hydrochloric and sulphuric acids, and readily in nitric acid. Like iron, it is magnetic. Alloys. — Alloys of nickel are extensively used. Argen- tan or German silver consists of copper, zinc, and nickel. Various nickel alloys are used for making coins. The 5-cent pieces in the United States are made of an alloy consisting of 25 per cent nickel and 75 per cent copper. In Switzerland, and Belgium also, nickel coins are used. Other Applications of Nickel. — Besides as a constituent of important alloys, nickel is extensively used at present in nickel-plating. Iron objects are covered with a thin layer of the metal for the purpose of protecting them from rusting. The plating is accomplished as silver-plating and copper-plating are — by means of electrolysis, a bath of nickel-ammonium sulphate being used. Nickelous Hydroxide, Ni(OH) 2 , is formed when a nickel salt is treated with a soluble hydroxide, and is a green insoluble substance. When heated it is converted into the green oxide, NiO. Nickelic Hydroxide, Ni(OH) 3 , is precipitated as a black powder when a solution of a nickel salt is treated with sodium hypochlorite. Cyanides. — When potassium cyanide is added to a solu- tion of a nickel salt, nickel cyanide, Ni(CN) 2 , is precipi- tated as a greenish-white substance. With an excess of potassium cyanide this forms the salt, Ni(CN) 2 .2KCN, which, owing to the fact that nickelous salts are not con- verted into nickelic salts by oxidation, does not undergo change when boiled with potassium cyanide. When hydrochloric acid is added to a solution of the double REACTIONS OF COBALT AND NICKEL. 639 cyanide, nickelous cyanide is precipitated. If boiled with precipitated mercuric oxide the double cyanide is decom- posed and nickel oxide is thrown down. Nickel Carbonyl, Ni(CO) 4 . — This interesting compound is formed when finely- divided nickel, such as is obtained by reducing nickel oxide by hydrogen at about 400°, is allowed to cool in a slow current of carbon monoxide. A gas is formed which can easily be condensed, its boiling- point being 43°. At — 25° it solidifies, forming needle- shaped crystals. When the gas is passed through a heated tube pure nickel is deposited. Advantage of this fact is taken for the purpose of preparing pure nickel on the large scale. Cobalt does not form a compound of this kind. Reactions of Cobalt and Nickel which are of Special Value in Chemical Analysis. — The reactions with the soluble hydroxides have been explained. With ammonium sulphide both give black sulphides, which are not easily dissolved by dilute hydrochloric acid. From solutions of the acetates hydrogen sulphide precipitates the sulphides. Nickel sulphide is slightly soluble in ammonium sulphide, and the solution has a brownish-yellow color. The action of the hypochlorites upon solutions of nickel and cobalt salts has been explained above. The reactions with potassium cyanide have also been explained. These furnish a good method for separating the two metals. When a solution of potassium nitrite is added to a solu- tion of a cobalt salt containing free acetic acid or nitric acid, a precipitate of cobaltic potassium nitrite is formed. This is a compound of cobaltic nitrite, Co(N0 2 ) s , and potassium nitrite, of the composition Co(N0 2 ) 3 .3KN0 2 . The formation involves oxidation of the cobaltous salt, and this is effected by some of the nitrogen trioxide which is set free. Thus with the chloride the action may be repre- sented as follows : CoCl, + 7KNO, + 2C 2 H 4 2 = 2KC1 + Co(N0 2 ) 3 .3KN0 2 + 2KC 2 H 3 2 + H 2 + NO. Nickel does not form a similar compound of a nickelic salt, 640 COLLEGE CHEMISTRY. but simply forms a double nitrite, containing the nickelous salt Ni(N0 2 ) 2 .4K]Sr0 2 . Cobalt compounds color the bead of microcosmic salt blue both in the reducing and oxidizing flame. Nickel colors it reddish brown in the oxidizing flame when hot, and pale yellow when cold. In the reducing flame it is gray. EXPERIMENTS. Ikon and its Compounds. Experiment 239. — Make ferric chloride by heating the purest iron wire in a current of chlorine. Also make a solution by dis- solving iron in hydrochloric acid, and oxidizing the solution with nitric acid. Experiment 240. — Make ferrous sulphate by dissolving iron in dilute sulphuric acid, and evaporating to crystallization. Dis- solve equivalent quantities of ferrous sulphate and ammonium sulphate, and evaporate to crystallization. Experiment 241.— Dissolve ferric hydroxide in sulphuric acid, and evaporate to dryness. CHAPTER XXXIV. ELEMENTS OF FAMILY VIII, SUB-GROUP B: RUTHENIUM.— RHODIUM.— PALLADIUM. . ELEMENTS OF FAMILY VIII, SUB-GROUP C: OSMIUM.— IRIDIUM— PLATINUM. General. — Comparing the members of the three sub- groups of Family VIII with reference to their atomic weights and specific gravities, we have the following remarkable table: Fe Co Ni At. Wt. 56 At. Wt. 59 At. Wt. 58.7 Sp. Gr. 7.8 Sp. Gr. 8.5 Sp. Gr. 8.8 Eu Eh Pd At. Wt. 101.7 At. Wt. 103 At. Wt. 106 Sp. Gr. 12.26 Sp. Gr. 12.1 Sp. Gr. 11.5 Os Ir Pt At. Wt. 191 At. Wt. 193 At. Wt. 194.8 Sp. Gr. 22.48 Sp. Gr. 22.42 Sp. Gr. 21.50 It will be observed that the atomic weights and specific gravities of the members of each sub-group are approxi- mately the same. But just as there is a gradual change in the chemical conduct from iron to nickel in the iron group, so a similar gradation of properties is observed in the other two groups. As far as the variety of com- pounds which they form is concerned, ruthenium and osmium are more like iron than they are like rhodium and iridium. Further, rhodium and iridium resemble each 641 642 COLLEGE CHEMISTRY. other, as regards the variety of their compounds, more closely than they resemble palladium and platinum, and a similar resemblance is noticed between palladium and platinum. A full discussion of these relationships would lead too far. The Platinum Metals, The six elements of Sub-Groups B and C, Family VIII, are generally grouped together and spoken of as the platinum metals. They occur together in nature, and almost always in alloys, into the composition of which all enter. The chief constituent is platinum, which is present to the extent of 50 to 80 per cent, and oyer. The alloys occur in only a few localities, in the Ural Mountains, in California, Australia, Borneo, and a few other places, and form small pieces which are mixed with sand and earth. They generally contain also gold, iron, and copper. Pal- ladium occurs, further, in a gold ore which is found in Brazil. Metallurgy. — The process for obtaining the metals from the ores is based mainly upon the following facts: (1) Gold is soluble in dilute aqua regia, while platinum requires concentrated aqua regia; (2) platinic chloride, Pt01 4 , and iridium chloride, Ir01 4 , form, with ammonium chloride, difficultly soluble compounds of the formulas (NH 4 ) 2 PtCl 6 (PtCl 4 .2NH 4 Cl) and (NH 4 ) 2 IrCl 6 (IrCl 4 .2Nn 4 Cl). When these compounds are ignited, they are completely decomposed, and the metals are left behind. When, therefore, platinum-ore has been freed as far as possible from sand and earth, it is first treated with dilute aqua regia, which removes the gold, and then with concentrated aqua regia, which dissolves- the platinum together with a little iridium, leaving an alloy of iridium and osmium. When the solution thus obtained is treated with am- monium chloride, both metals are precipitated; and when the precipitate is ignited, both metals are left behind in the form of a spongy mass. This consists, however, almost RUTHENIUM— OSMIUM. 643 wholly of platinum, the amount of iridium being very small. Ruthenium, Ku (At. Wt. 101.7). Preparation. — Ruthenium is, obtained from the residue which is left undissolved when platinum-ore is treated with concentrated nitro-hydrochloric acid. Properties. — When heated in oxygen it burns and forms the oxide, Ru0 2 . It is insoluble in the strong acids, and even in nitro-hydrochloric acid it is almost insoluble. Owing to its power to form salts of ruthenious acid, it is dissolved when heated with potassium hydroxide and an oxidizing agent, such as saltpetre or potassium chlorate, and afterwards treated with water. Osmium, Os (At. Wt. 191). Preparation. — As stated above, this element is left un- dissolved in the form of an alloy with iridium when platinum-ore is treated with concentrated nitro-hydro- chloric acid. In order to separate it from the iridium, advantage is taken of the fact that it forms a volatile peroxide, Os0 4 , similar to that formed by ruthenium, while iridium does not. Properties. — The metal does not melt at the highest temperatures reached artificially. It has the highest specific gravity of all known substances; is easily oxidized when in finely divided condition; and is converted either by the oxygen of the air or by nitric acid into osmium peroxide, Os0 4 . The metal as well as the oxides forms the peroxide, 0s0 4 , when heated in the air. This is also formed by treating a heated mixture of sodium chloride and the alloy of osmium and iridium with chlorine and water vapor. It is commonly called osmic acid, though its acid properties are very weak. Like ruthenium peroxide it is volatile. It sublimes in colorless, lustrous needles, and boils without decomposition at a temperature 644 ' COLLEGE CHEMISTRY. a little above 100°. It lias an intense odor similar to that of chlorine, and its vapor attacks ^he eyes and respiratory organs somewhat as chlorine does. It dissolves slowly in water, and reducing agents precipitate the metal from the solution. A solution of osmic acid is used in microscopic work. When injected into the tissues, the parts are hardened and colored. Rhodium, Eh (At. Wt. 103). Rhodium has no acid properties, and does not form a peroxide corresponding to those of ruthenium and osmium. On the other hand, its oxide, Rh 2 3 , is basic. The chlo- ride RhCl 3 is readily formed. It is doubtful whether the di- and tetrachlorides have been made. Iridium, Ir (At. Wt. 193). Preparation. — The extraction of iridium with platinum and with osmium from platinum-ore was referred to above. In order to separate it from platinum, advantage is taken of the fact that it forms a trichloride, IrCl 3 , which with ammonium chloride gives an easily soluble double chloride. Properties. — Iridium has a grayish-white color, and resembles polished steel. Its specific gravity is nearly the same as that of osmium, being 22.42. It is harder and more brittle than platinum; melts at a higher temperature; and unless it is finely divided it is not dissolved by nitro- hydrochloric acid. When heated with potassium hydroxide and saltpetre it is converted into the oxide. Palladium, Pd (At. Wt. 106). Preparation. — The chief source of palladium is a Brazilian gold-ore. From this ore the metal can be obtained by various methods, one of which consists in melting it together with silver, and then treating it with nitric acid, when the silver and palladium dissolve, and the gold remains undissolved. The silver is precipitated PALLADIUM— PLATINUM. 645 as chloride and the palladium as the cyanide, and when the latter is ignited it is decomposed, leaving palladium. Properties. — Palladium resembles iridium and platinum in appearance. Its specific gravity is only about half as great as that of platinum, being 11.5; it is more easily fusible than platinum, and dissolves in nitric acid and in hot concentrated sulphuric acid. The property of pal- ladium which has perhaps attracted most attention is its power to absorb hydrogen, and form Palladium-Hydrogen — The formation of this compound was referred to under Hydrogen (which see). The com- bination takes place even at the ordinary temperature, but best at 100°. If the metal is brought into hydrogen at this temperature, it absorbs more than 900 times its volume, forming an alloy of the composition Pd 2 H. This alloy has a greater volume and lower specific gravity than the palladium from which it is formed. At 130° it begins to decompose under the atmospheric pressure, but con- tinued heating at a red heat is necessary to decompose it completely. If allowed to lie in contact with the air the hydrogen is oxidized to water. Palladium-hydrogen acts as a strong reducing agent, the hydrogen which it gives up being apparently in the nascent or atomic condition. Platinum, Pt (At. Wt. 194.8). Preparation. — A general idea of the method of pro- cedure in extracting platinum from its ores was given on p. 642. Thus prepared, however, it always contains iridium, and for some purposes for which platinum is used this is objectionable. In order to purify the metal advan- tage is taken of the fact that iridium chloride can be con- verted into a trichloride, which with ammonium chloride forms an easily soluble double salt. The metal as ob- tained by igniting ammonium platinic chloride forms a gray spongy mass known as spongy platinum. When a solution of platinous chloride is boiled with potassium hydroxide, and alcohol gradually added, the salt is re- 646 COLLEGE CHEMISTRY. duced, and the platinum is precipitated as an extremely fine powder, known as platinum Hack. When spongy platinum and platinum-black are • heated to fusion by the oxyhydrogen flame they are converted into the compact variety. Properties. — Platinum is a grayish-white metal resem- bling polished steel; it can be drawn out into very fine wire; it melts in the flame of the oxyhydrogen blowpipe, and when heated above its melting-point it is volatile; its specific gravity is 21.5. At white heat it can be welded. It is not dissolved by nitric acid, hydrochloric acid, or sulphuric acid, but it dissolves in nitro-hydrochloric acid, forming the acid, E^PtCl,.. Fusing akalies, and particu- larly a mixture of caustic potash and saltpetre, act upon it ; but the alkaline carbonates do not. In contact with red-hot charcoal and silicon dioxide a compound of silicon and platinum is formed. Finely divided platinum has to a remarkable extent the power of condensing gases upon its surface. It absorbs, for example, 200 times its own volume of oxygen, and other gases in a similar way. The oxygen thus absorbed is in an active condition, and if oxidizable substances are brought in contact with it they are easily oxidized. Thus when a current of hydrogen is allowed to flow against a piece of spongy platinum it takes fire, owing to the presence of the condensed oxygen in the pores of the platinum. Similarly, when sulphur dioxide and oxygen are allowed to flow together over spongy platinum, or even the compact metal, the two gases unite to form sulphur trioxide. Applications of Platinum. — The metal is of great value to the chemist on account of its power to resist high tem- peratures and the action of most chemical substances. It is used in the laboratory in the form of wire, foil, cruci- bles, evaporating-dishes, tubes,* etc., etc. From what was said above it cannot be used with alkalies and saltpetre, nor with nitro-hydrochloric acid. Platinum vessels, further, should not be placed upon red-hot charcoal. Metallic salts which are easily reduced, such as those of ALLOYS OF PLATINUM. 647 antimony and bismuth, should not be heated in platinum vessels, as the reduced elements, like silicon, form alloys with the platinum, and these, as a rule, are easily fusible. In the concentration of sulphuric acid on the large scale platinum stills are used. The price of platinum is not as high as that of gold, but much higher than that of silver. Alloys of Platinum. — The only alloy of platinum which is of any special importance is that which it forms with iridium. A small percentage of iridium diminishes the malleability of platinum very markedly, and makes it brittle; it, however, increases its resistance to the action of reagents. An alloy of 90 per cent platinum and 10 per cent iridium has been adopted by the French Government as the best material from which to make normal meters. This alloy is very hard, as elastic as steel, more difficultly fusible than platinum, entirely unchangeable in the air, and is capable of a high polish. Chlorides. — Like palladium, platinum forms two chlo- rides, platinous chloride, PtCl 2 , and platinic chloride, PtCl 4 . Platinous chloride is formed by passing chlorine over platinum sponge heated to 240° to 250°. Platinic chloride is formed by heating chlorplatinic acid, H 2 PtCl 6 , in a current of chlorine to 360°. Platinic chloride is easily soluble in water. Chlorplatinic Acid, H 2 PtCl 6 , is formed by direct union of platinic chloride with hydrochloric acid, and by dissolv- ing platinum in aqua regia and evaporating the solution. It crystallizes with six molecules of water, and forms a series of salts called the chlorplatinates, to which reference has already been made. Those most commonly met with in the laboratory are the potassium salt, K 2 PtCl 6 , or PtCl t .2KCl, and the ammonium salt, (NH 4 ) 3 PtCl 6 , or PtCl 4 .2NH 4 Cl, both of which are difficultly soluble in water, and are therefore precipitated when chlorplatinic acid (platinic chloride) is added to solutions containing potassium or ammonium chloride. The sodium salt is easily soluble in water. Chlorplatinic acid appears to be analogous to fluosilicic acid, H 2 SiP 6 . 6 48 COLLEGE CHEMISTRY. Sulphides. — There are two sulphides of platinum which are analogous to the two oxides, PtO and Pt0 2 . These are platinous sulphide, PtS, and platinic sulphide, PtS . They are black insoluble compounds, which are precipi- tated when hydrogen sulphide or soluble sulphides are added to solutions of platinous and platinic chlorides. EXPERIMENT. Platinum. Experiment 242. — Prepare a solution of chlorplatinic acid as follows : Heat platinum in a flask with concentrated nitric acid, adding from time to time a few drops of hydrochloric^ acid. After the metal is dissolved evaporate to dryness on a water-bath. Dissolve in water and filter. Chemical Analysis. At the end of each chapter treating of the metallic or base- forming elements there is given a list of such reactions as are of special value for analytical purposes, together with such explana- tory statements as seem called for. The student who is engaged in analytical work will generally find these explanations sufficient to enable him to keep his ideas clear in regard to the reactions with which he is dealing, provided, at the same time, he carefully studies the chapter to which the explanations form an appendix. As an introduction to analytical work, a general study of chem- ical reactions is necessary, and the fuller this is the better. There are many small books in existence in which good directions are given for work of this kind. The student is, however, advised to supplement the book he may be using by such experiments as may suggest themselves on reading the corresponding chapters in this book. The directions there found will generally be quite sufficient for the purpose, and it is therefore not considered nec- essary to give more specific directions in this place. In general, the more the student occupies himself in the laboratory with chemical substances the more rapidly will his chemical ideas grow. But it is necessary that he should avoid working by '■ rule of thumb," and for this purpose constant reference to some larger text-book in which the relations between the substances and the reactions he is dealing with are discussed in a broad way is of the highest importance. CHAPTER XXXV. SOME FAMILIAR COMPOUNDS OF CARBON. Organic Chemistry.— When the compounds that are obtained from plants and animals were first studied, they were supposed to be entirely different from the compounds obtained from the inorganic, er mineral, constituents of the earth. The former were called organic compounds because they were obtained from organized things; while the latter were called inorganic compounds. Organic compounds were the subject of Organic Chemistry, and inorganic compounds formed the subject of Inorganic Chemistry. These names are still in use, though they have lost their original meaning. Organic Chemistry now means only the Chemistry of the Compounds of Carlon. Occurrence of the Compounds of Carbon. — The com- pound of carbon that occurs most widely distributed in nature is carbon dioxide. This, as has been pointed out, is the starting-point of all life on the globe. All living things are formed from it either directly or indirectly. Attention has been called to the fact that starch and cellulose are the principal compounds found in plants, and that fats, albumin, and fibrin are the most common sub- stances found in animals. Formation of Hydrocarbons. — Certain natural processes which are not thoroughly understood have given rise to the formation of a complex mixture of organic com- pounds, principally hydrocarbons, in petroleum. Distillation of Coal. — The destructive distillation of coal for the purpose of making illuminating-gas, and the formation of coal-tar, have already been referred to. 649 6S° COLLEGE CHEMISTRY. Coal-tar is one of the most important sources of compounds of carbon. The hydrocarbons benzene, C 6 H 6 , toluene, C,H 8 , xylene, O 8 H 10 , naphthalene, C 10 H 8 , anthracene, 14 H 10 , etc., are obtained from this source. Distillation of Wood. — Wood is heated in closed vessels mostly for the purpose of making charcoal, as already explained. Among the products obtained from this source are wood-spirit, or methyl alcohol, and pyroligneous acid, or acetic acid. Large quantities of acetic acid are prepared in this way. Distillation of Bones — In order to make bone-black, bones are subjected to destructive distillation. The oil which passes over is collected and known as bone-oil. This is the source of a large number of compounds which are of special interest on account of their connection with the valuable alkaloids quinine, morphine, etc. Fermentation. — A number of the most important com pounds of carbon are formed by a process known as fermen- tation. This is a general term meaning any process in which a chemical change is effected by means of minute animal or vegetable organisms. The best-known example of fermentation is that of sugar, which gives rise to the formation of ordinary alcohol. Classes of Compounds of Carbon. — The chief classes of these compounds are the hydrocarbons, the alcohols, the aldehydes, the acids, the ethers, and the ethereal salts. First a few of the best-known examples of each of these classes will be taken up, and afterwards some other familiar compounds which do not belong to any one of these classes. HTDKOCAEBONS. Compounds of Carbon and Hydrogen. — It is not an easy matter to effect combination between carbon and hydrogen in the laboratory except in a few simple cases. In nature processes are in operation which give rise to the formation of a large number of compounds containing these elements; and, further, in the manufacture of iiluminating-gas from PETROLEUM. 651 coal the conditions are such as to cause the combination of carbon and hydrogen, several interesting compounds being thus formed. There are no other two elements which combine with each other in as many different pro- portions as carbon and hydrogen. The compounds thus formed are known as hydrocarbons. The number of hydrocarbons known is large, being somewhere near two hundred. Fortunately, investigation has shown that quite simple relations exist between these compounds; and hence, though the number is large, the study is not as difficult as might be expected. Petroleum is an oily liquid found in many places in the earth in large quantity, particularly in Pennsylvania and the Caucasus. In the earth it contains both gases and liquids. When it is brought into the air, the pressure being removed, the gases are given off. There are several gaseous hydrocarbons given off, and a large number of liquids left behind. Refining of Petroleum. — The vapors from petroleum when mixed with air are explosive, and the thicker liquids clog the lamps and wicks. Therefore these must be removed before the oil is fit for household use. This is done by (1) distilling, (2) washing with sulphuric acid, (3) washing with alkali, and (4) washing with water. The product thus prepared is called kerosene. In refining petroleum a number of products are obtained which cannot be used in lamps. Those which are lighter than kerosene, that is to say those which boil at a lower temperature, are known as gasoline, naphtha, benzine, etc. From the heavier portions, or those which boil at higher temperatures than kerosene, lubricating oils and paraffin are made. Each of these substances is a mixture of several chemical compounds. Hydrocarbons contained in Petroleum. — The simplest hydrocarbon contained in petroleum is methane, or marsh- gas, CH 4 ; the next has the composition 2 H 6 , the next C S H , etc. It will be seen that these compounds bear a simple relation to one another, as far as composition is ^S 2 COLLEGE CHEMISTRY. concerned. They are the first members of a series the names and symbols of the first eight members of which are given below: CH 4 , Methane, or Marsh- ■gas; CA> Ethane ; C 3 H 8> Propane; CA 0J Butane; C 5 H 12> Pentane ; °6 H H' Hexane; 7 H 16> Heptane ; C 8 H 18> Octane. Homology. — The first member of the series differs from the second by CH 2 ; there is also this same difference, in general, between any two consecutive members of the series. This relation is known as homology, and such a series as an homologous series. Carbon is distinguished from all other elements by its power to form homologous series. The Ethylene 'Series of Hydrocarbons. — Besides the series above mentioned, which is known as the marsh-gas series, there are other homologous series of hydrocarbons. There is one beginning with ethylene, C 2 H 4 , examples of which are Ethylene, C 2 H 4 ; Propylene, C 3 H 6 ; Butylene, C 4 H 8 . The Acetylene Series. — There is a series beginning with acetylene, examples of which are Acetylene, C 2 H 2 ; Allylene, C 3 H 4 . The Benzene Series. — Another series begins with ben- zene, C 6 H 6 . Some of the members of this series are Benzene, 6 H 6 ; Toluene, C,H 8 ; Xylene, C 8 H 10 . MARSH-GAS. 653 Marsh-gas,. Methane, Fire-damp, CH + . — Marsh-gas is found in nature in petroleum, and is given off when the oil is taken out of the earth, and the pressure is removed. It is formed, as the name implies, in marshes, as the product of a reducing process. Vegetable matter is com- posed of carbon, hydrogen, and oxygen. When it under- goes decomposition in the air in a free supply of oxygen, the final products formed are carbon dioxide and water. When the decomposition takes place without access of oxygen, as under water, marsh-gas, which is a reduction- product, is formed. The gas can be made in the laboratory by treating aluminium carbide, C 3 Al 4 , with water : C 3 A1 4 + 12H 2 = 3CH 4 + 4Al(OH),. It can also he made by passing a mixture of hydrogen sul- phide and the vapor of carbon bisulphide over heated copper. Marsh-gas is met with in coal-mines, and is known to the miners as fire-damp, "damp" being the general name applied to a gas, and the name fire-damp meaning a gas that burns. To prepare it in the laboratory, it is most convenient to heat a mixture of sodium acetate and quicklime. The change which takes place will be readily understood by considering it as a simple decomposition of acetic acid. Acetic acid has the formula C 2 H 4 0„. When heated alone, it boils and does not suffer decomposition. If it is converted into a salt, and heated in the presence of a base, it breaks down into marsh-gas and carbon dioxide: C 2 H,0 2 = CH 4 + C0 2 . The carbon dioxide, which forms salts with bases, does not pass off, but remains behind in the form of a salt of car- bonic acid. Marsh-gas is a colorless, transparent, tasteless, inodorous. rras. It is slightly soluble in water. It burns, forming carbon dioxide and water. When mixed with air, the 654 COLLEGE CHEMISTRY. mixture explodes if a flame or spark comes in contact with it. This is one of the causes of the explosions which so frequently occur in coal-mines. To prevent these explo- sions a special lamp was invented by Sir Humphry Davy, which is known as Davy's safety-lamp (p. 374). Substitution-products of the Hydrocarbons. — Marsh-gas and other hydrocarbons undergo change when treated with chlorine and bromine. The change consists in the substi- tution of one or more atoms of chlorine or of bromine for the same number of atoms of hydrogen. In the case of marsh-gas and chlorine the possible changes are repre- sented as below: CH 4 + Cl 2 = CH,C1 + HC1; CH 3 C1 + Cl 2 = CH 2 C1 2 + HC1; CH 2 C1 2 + Cl 2 = CHC1 3 + HC1; OHCl a + Cl 2 = CC1 4 + HC1. All the products represented are known. Chloroform, CHC1 3 . — Chloroform can be made as above indicated, but it is made on the large scale by treating alcohol (which see) or acetone (which see) with bleaching- powder. It is a heavy liquid with an ethereal odor and a somewhat sweet taste. It is one of the most valuable anaesthetics, though there is some danger attending its use. Iodoform, CHI 3 . — This compound, like chloroform, is a substitution-product of marsh-gas. It is made by bring- ing together alcohol, an alkali, and iodine. It is a solid substance, soluble in alcohol and ether, but insoluble in water. It crystallizes in six-sided yellow plates. It is extensively used as a dressing for wounds in surgery. Ethylene, Olefiant Gas, C 2 H 4 . — This hydrocarbon is formed by heating a mixture of ordinary alcohol and con- centrated sulphuric acid. The reaction is represented thus: C 2 H 6 = H 2 + C 2 H 4 . Alcohol. Ethylene. Ethylene is a colorless gas, which can be condensed to ALCOHOLS. 655 a liquid. It burns with a luminous flame. With oxygen it forms an explosive mixture. Acetylene, C 2 H 2 . — Acetylene is formed when a current of hydrogen is passed between carbon poles, which are in- candescent in consequence of the passage of a powerful electric current. In this case carbon and hydrogen com- bine directly. It is formed also when the flame of an ordinary laboratory gas-burner (Bunsen burner) " strikes back," or burns at the base without a free supply of air. It is now made on the large scale by treating calcium car- bide, CaC 2 , with water: CaC 2 + H 2 = C 2 H 2 + CaO. Its odor is unpleasant.. It burns with a luminous, smoky flame. ALCOHOLS. Methyl Alcohol, Wood-spirit, CH 4 0. — This is formed in the distillation of wood together with other products. It has, when pure, a pleasant odor and taste, and acts upon the animal system very much as ordinary alcohol does. It burns without giving light or smoke, and can therefore be used in lamps for heating-purposes as ordinary alcohol is. It is used in the manufacture of varnishes. Ethyl Alcohol, Spirits of Wine, C 2 H 6 0.— This well- known substance is formed by the fermentation of grape- sugar or glucose. What Change takes Place in the Sugar ? — If the solu- tion in the flask is examined carefully it will be found to contain alcohol and no sugar. Grape-sugar has the com- position expressed by the formula C 6 H 12 6 . By fermenta- tion it is decomposed, forming alcohol, C 2 H 6 0, and carbon dioxide, C0 2 , thus: C 6 H 12 6 = 2C 2 H 6 + 2C0 2 . What Causes the Change ? — It has been found that the change of grape-sugar is caused by small organized bodies 6*6 COLLEGE CHEMISTRY. rhich grow in the solution. These bodies are contained in ordinary yeast. Germs in the Air. — When fruit-juices that contain sugar are exposed to the air they undergo fermentation with- out the addition of yeast. This is due to the fact that the germs or seeds of the bodies which cause fermentation are everywhere floating in the air. Hence when a liquid in which these seeds can grow is exposed to the air, the bodies are formed and fermentation takes place. Different Kinds of Fermentation. — The fermentation " which yields alcohol is only one of many kinds. Among the others are : (1) lactic-acid fermentation, which takes place in the souring of milk; and (2) acetic-acid fermenta- tion, which causes the transformation of alcohol into acetic acid. The latter ferment is contained in "mother of vinegar." Distillation of Fermented Liquids, — In order to get the alcohol from liquids which have undergone fermentation they must be distilled. For this purpose very perfect forms of stills have been devised, so that the alcohol passes over nearly free from other substances. Usually it con- tains impurities known as fusel oil. Properties of Alcohol. — Pure ethyl alcohol has a peculiar, pleasant odor. It remains liquid at very low temperatures, but has been converted into a solid at a temperature of — 130.5°. It burns with a flame which does not deposit soot, and was hence formerly much used in laboratories for heating purposes, and is still used where gas cannot be obtained. Its effects upon the human system are well known. It intoxicates when taken in dilute form, while-- in large doses it is poisonous. It lowers the temperature of the body when taken internally, although it causes a sensation of warmth. Uses of Alcohol. — Alcohol is the principal solvent for organic substances. It is hence extensively used in the arts, as in the manufacture of varnishes, perfumes, and tinctures of drugs. Most beverages in use owe their in- toxicating-power to the presence of alcohol. The milder GLYCERIN- ALDEHYDES. 657 forms of beer contain from 2 to 3 per cent; light wines about 8 per cent; while whiskey, brandy, etc., sometimes contain as much as 60 to 75 per cent. Glycerin, C 3 H e 3 . — Glycerin is an alcohol which occurs very widely distributed as a constituent of fats. The relation it bears to the fats will be explained when the acids which enter into the fats are taken up. It is obtained from the fats by boiling them with an alkali like caustic soda or caustic potash, or by heating with steam. Properties. — Glycerin is a thick, colorless liquid with a sweetish taste. It attracts moisture from the air, and is hence used to keep surfaces moist. ALDEHYDES. Acetic Aldehyde, Ordinary Aldehyde, C 2 H 4 0. — This compound is formed by oxidizing ordinary alcohol, the change being represented by this equation : C 2 H 6 + = C 2 H 4 + H 2 0. Aldehyde is a volatile liquid with a characteristic pene- trating odor. When left to itself, and especially when treated with a number of other things, it is converted into another substance of the same composition. This is called paraldehyde. A determination of the molecular weight of the substance by the method of Avogadro has shown that it must be represented by the formula C 6 H 12 3 . The change from aldehyde to paraldehyde must, therefore, be represented thus: 3C 2 H 4 = C 6 H 12 3 . Paraldehyde is used in medicine. Chloral, G 2 C1 3 H0, is a compound formed by the action of chlorine on alcohol. It is related to aldehyde, as chloroform is related to marsh-gas, that is to say, it is a trichlorine substitution-product. It is a colorless liquid. With water it forms a crystallized compound, chloral hydrate, 2 C1 3 H0 + H 2 0, which is easily soluble in water, and crystallizes from the solution in colorless prisms. 65 8 COLLEGE CHEMISTRY. Taken internally in doses of from 1.5 to 5 grams, it produces sleep. In larger doses it acts as an anaesthetic. ACID8. Formic Acid, CH 2 2 . — This acid occurs in nature in red ants, in stinging-nettles, in the shoots of some of the varieties of pine, and elsewhere. It is a colorless liquid. Dropped on the skin, it causes extreme pain and produces blisters. Acetic Acid, C a H 4 2 . — This is the acid contained in vinegar, and the value of vinegar is due to its presence. It is formed from alcoholic liquids by exposing them to the air, in consequence of the presence of a microscopic organism which is contained in what is commonly known as "mother of vinegar." The formation of acetic acid from alcohol is due to the action of oxygen as represented in the equation C 2 H 6 O + O 2 = 2 H 4 O 2 + H 2 0. Alcohol. Acetic acid. But oxygen alone does not effect the change. When the ferment is present the oxidation takes place. Acetic acid is also obtained by distilling wood. Hence the names pyroligneous acid and wood-vinegar. Properties. — Acetic acid is a clear, colorless liquid. It has a very penetrating, pleasant, acid odor, and a sharp taste. The pure substance acts upon the skin like formic acid, causing pain and raising blisters. Uses. — Acetic acid is extensively used, chiefly in the dilute form known as vinegar. It is used in calico-print- ing in the form of iron and aluminium salts. With iron it gives hydrogen, which is needed in the manufacture of certain compounds used in making dyes. Salts of Acetic Acid. — The best-known salts of acetic acid are lead acetate, Pb(C 2 H 3 2 ) 2 , commonly called sugar of lead; and copper acetate, Cu(C 2 H,0 2 ) 2 , a variety of which is known as verdigris. Fatty Acids. — Formic and acetic, acids are the first members of an homologous series (see p. 652). Some of FATTY ACIDS. 659 the more important members are named in the following table : Formic acid CH a O a . Acetic " C a H 4 O a . Propionic " C,H,O a . Butyric " C^H.O,. Palmitic " ;... C 18 H sa O a . Stearic " C„H. s O a . They are called fatty acids for the reason that many of them are obtained from fats. Butyric acid, C 4 H 8 2 , is of special interest because it is obtained from butter by boiling with caustic potash. It occurs also in many other fats. There is a butyric-acid ferment contained in putrid cheese which has the power of converting sugar into butyric acid. Palmitic acid, C 16 H S2 2 , is obtained from many fats, but palm-oil is especially rich in it. Stearic acid, 18 H S6 2 , is the acid contained in the fat known as stearin. The so-called "stearin candles" are made of a mixture of palmitic and stearic acids. Soaps. — Soaps are the alkali salts of the acids contained in fats, especially of palmitic and stearic acids. Fats are compounds of these acids with glycerin. When the fats are boiled with an alkali, as caustic soda, the correspond- ing salts of the acids are formed, while the glycerin is set free. The potassium and sodium salts of palmitic and stearic acids are the soaps. Use of Soap. — The cleansing power of soap depends upon the fact that it dissolves the oily film on the surface of the skin and thus facilitates the removal of the foreign sub- stances commonly known as dirt. Action of Soap on Hard Waters. — As has been explained, a hard water is one that contains salts in solution. Tem- porary hardness is that which is caused by calcium car- bonate held in solution in the water by carbon dioxide. Permanent hardness is caused by calcium sulphate or magnesium salts. The calcium and magnesium salts of palmitic and stearic acids are insoluble in water. There- 660 COLLEGE CHEMISTRY. fore, when soap is added to a hard water these insoluble salts are precipitated and give the water a hard feeling. In attempting to wash the hands with soap in a hard water they become covered with a thin layer of the insoluble salts which prevents them from rubbing freely over each other, and makes them feel sticky. Before the soap can do any good all the lime-salt must be precipitated. The action in the case of temporary hardness is represented by the equation 2NaC 16 H 31 2 + CaC0 3 = Ca(C 16 H 31 2 ) 2 + Na 2 C0 3 . Soap. Calcium palmitate. In the case of permanent hardness it is represented by the equation 2NaC 16 H 31 2 + CaS0 4 = Oa(0 I ,H, 1 1 ), + Na 2 S0 4 . Relations of the Soap Industry to other Industries A great chemist and philosopher has said that the quantity of soap used in a country is a measure of the civilization of that country. Certain it is that soap is only used by civilized people, and that by them it is used in enormous quantities. In many farm-houses a primitive method for the manufacture of soap is practised, consisting in treating refuse fats with the lye extracted from wood-ashes. A soft soapy mass is thus obtained known as "soft-soap." Fats form the starting-point in the manufacture of all soap. These are generally treated with caustic soda. Caustic soda is all made from sodium carbonate by the action of lime; and, as has been seen, sodium carbonate is made from common salt mostly by the Le Blanc process, which re- quires sulphuric acid. Thus the manufacture of sulphuric acid and sodium carbonate is intimately related to the manufacture of soap. Oxalic Acid, C 2 H 2 4 . — This acid occurs very widely dis- tributed in nature, as in the sorrels, which owe their acid taste to the presence of acid* potassium oxalate, KC 2 H0 4 ; and as the ammonium salt in guano. It is probably one of the first substances formed from carbon dioxide in the ETHERS. 66 1 plant. It is manufactured by heating wood shavings or sawdust with caustic soda and caustic potash. Oxalic acid is an active poison. It is used in calico-printing, and in cleaning brass and copper surfaces. Lactic Acid, C 3 H 6 3 . — Lactic acid is made by the fer- mentation of sugar by means of the lactic-acid ferment. The reaction effected by the ferment is represented by the equation C 6 H 12 6 = 2C 3 H 6 O s . Malic Acid, C 4 H 6 5 . — This acid is very widely distributed in the vegetable kingdom, as in apples, cherries, etc. Tartaric Acid, C 4 H 6 6 . — Tartaric acid occurs very widely distributed in fruits, sometimes uncombined, sometimes in the form of the potassium or calcium salt; as, for example, in grapes, berries of the mountain-ash, potatoes, cucum- bers, etc., etc. It is prepared from "cream of tartar." This is acid potassium tartrate, which is formed when grape-juice ferments. Citric Acid, C 6 H 8 7 . — Citric acid, like malic and tartaric acids, is very widely distributed in nature in many varieties of fruit, especially in lemons. It is also found in currants, whortleberries, raspberries, gooseberries, etc., etc. It is prepared from lemon- juice: 100 parts of lemons yield 5£ parts of the acid. It is a solid, crystallized substance, soluble in water. It is frequently used for the purpose of making lemonade without lemons, and there is no objec- tion to its use for this purpose. ETHERS, Ether, C 4 H 10 0. — Ordinary ether is the best-known repre- sentative of the class of compounds called ethers. It is formed from ordinary alcohol by treating it with sulphuric acid and distilling. The result of the action which takes place is represented by the equation 2C 2 H 6 = C 4 H 10 O + H 2 0. Alcohol. Ether. 662 COLLEGE CHEMISTRY. Ether is a liquid which boils at a low temperature and takes fire and burns readily. Inhaled it produces insensi- bility to pain. It is therefore called an ancesthetic. ETHEREAL SALTS. Action of Acids upon Alcohols — When an acid acts upon an alcohol it is neutralized, though not as readily as when it acts upon a base. The product is a substance which resembles a salt and is called an ethereal salt. Thus when nitric acid acts upon alcohol this reaction takes place: C 2 H 6 + HNO s = C 2 H 6 N0 3 + H 2 0. The product C 2 H 5 N0 3 , called ethyl nitrate, is an ethereal salt. The alcohol acts as if it were a substance like caustic potash and made up thus: C 2 H 5 OH. The resemblance between its action and that of caustic potash is shown by the equations KOH + HNO s = KN0 3 + H 2 0, and C 2 H 6 OH + HN0 3 = C 2 H 5 N0 3 + H 2 0. Saponification. — When an ethereal salt is boiled with a caustic alkali it is decomposed, the products being an alcohol and an alkali salt. Thus when ethyl nitrate is boiled with caustic potash, potassium nitrate and alcohol are formed : C 2 H 6 N0 3 + KOH = C 2 H 5 OH + KN0 3 . This process is called saponification, because the most important example is furnished by soap-making. Fats. — The fats are ethereal salts in the formation of which glycerin, as the alcohol, and three acids take part. The three acids are palmitic and stearic acids, already mentioned, and oleic acid, C 18 H 34 2 . Although the com- position of these substances is comparatively Complex, the way they act upon one another is simple, and is the same as the action of nitric apid upon alcohol in forming ethyl nitrate. The fats, then, are the palmitate, stearate, and oleate of glyceryl, which bears to glycerin very much the ETHEREAL SALTS, ETC. 663 same relation that ethyl, C 2 H 5 , bears to alcohol. When a fat is boiled with caustic soda, glycerin and the sodium salts of the acids contained in the fat are formed. Butter consists of ethereal salts of glycerin and several fatty acids, among which are palmitic, stearic, and butyric. Oleomargarin is an artificial butter made from other fats than that of milk. Ethereal Salts as Essences The ethereal salts generally have pleasant odors, and it is to their presence that many fruits owe their flavors. Some of the compounds are now made artificially and used instead of the fruit-extracts. Thus the ethyl salt of butyric acid is used under the name of essence of pineapples, and the amyl salt of valeric acid under the name of essence of apples. Nitroglycerin. — Among the more important ethereal salts of glycerin are the nitrates. Two of these are ro.N0 2 known, viz., the mono-nitrate, C S H 5 ■; OH , and the tri- (OH • nitrate, C,H 6 (O.NO,),, the latter being the chief con- stituent of nitroglycerin. Nitroglycerin is prepared by treating glycerin with a mixture of concentrated sulphuric and nitric acids. It is a pale yellow oil which is insoluble in water. At — 20° it crystallizes in needles. It explodes very violently by concussion. It may be burned in an open vessel, but if heated above 250° it explodes. Dynamite is infusorial earth* impregnated with nitro- glycerin. Nitroglycerin is the active constituent of a number of explosives. RELATIONS BETWEEN" THE COMPOUNDS CONSIDERED. Comparison of the Formulas. — On comparing the formu- las of the hydrocarbons of the marsh-gas series (see p. 652) with those of the simplest alcohols and the fatty acids, it will be seen that these compounds are all related in a * That is to say, earth made up of the microscopic flinty shells which constitute the fossil remains of certain minute and simple plants. 664 COLLEGE CHEMISTRY. simple way. Below are the formulas of a few of the hydrocarbons, alcohols, and acids: Hydrocarbons. Alcohols. Acids. CH 4 CH 4 0H.O, C 2 H 6 C 2 H 4 2 3 H 8 C s H 8 C,H 6 2 C 4 H 10 , etc. C 4 H 10 O, etc. „ etc. Each of these series is an homologous series. Alcohols. — Alcohols have been shown to be derived from the hydrocarbons by the replacement of one or more hydrogen atoms by oxygen and hydrogen, OH, called hydroxyl, or from water by replacing one of the hydrogen atoms of the water by a compound of carbon and hydro- gen. An alcohol, then, is a hydroxide, just as a metallic base is; only, instead of consisting of a metal in combina- tion with hydrogen and oxygen, it consists of a compound of carbon and hydrogen in combination with hydrogen and oxygen. Thus : Metallic Bases. Alcohols. K(OH) CH s (OH) Na(OH) C 2 H 6 (OH) More Complex Alcohols. — Just as lime is a more complex base than caustic potash, as shown by the formulas KOH and Ca0 2 H 2 or Ca(OH) 2 , so there are more complex alcohols than ordinary alcohol. A good example is fur- nished by glycerin, 3 H 8 5 , which has been shown to be a hydroxide corresponding to aluminium hydroxide, Al(OH) 3 , a fact which is represented by the formula C 3 H 6 (OH) s . It may be called glyceryl hydroxide, the complex, C 3 H 5 , being known as glyceryl. Radicals or Residues. — The compounds of hydrogen and carbon contained in the alcohols are called radicals or residues. We may say that an alcohol is water in which a radical has been substituted for half of the hydrogen. ACIDS. 665 HOH Water. HOH HOH HOH Water. C 2 H 5 OH Ordinary alcohol. ( OH 0,H 5 OH = C s H 8 3 . OH Glycerin. Acids — Just as the alcohols have been shown to be derived from water, so the organic acids have been shown to be derived from carbonic acid. Carbonic acid itself is not known. But the carbonates are de- rived from an acid of the formula H 2 CO s , or CO j qS. If, in this acid, a hydroxyl is replaced by a radical, as, for example, by ethyl, 2 H 5 , a substance of the formula CO j G ^ or C 3 H 6 2 is the result. If methyl, CH S , is ( OH introduced in place of ethyl, the product is CO ■! ^tt 3 or G s H 4 2 , which is acetic acid. In a similar way all the organic acids are derived from carbonic acid. EXPERIMENTS. Fermentation. Experiment 243. — Dissolve about 150 grams of commercial grape-sugar in !■£ litres of water in a flask. Connect the flask by Fig. 79. means of a bent glass tube with a cylinder or bottle containing clear lime-water. The vessel containing the lime water must be provided with a cork with two holes. Through one of these 666 COLLEGE CHEMISTRY. passes the tube from the fermentation-flask ; through the other a tube connected with a vessel containing solid caustic potash, the object of which is to prevent the air from acting upon the lime-water. The arrangement of the apparatus is shown in Fig. 79. Now add to the solution of grape-sugar a little fresh brewers' yeast ; close the connections and allow to stand. Soon an evolution of gas will begin, and, as this passes through the lime-water, a precipitate will be formed which can be shown to be calcium carbonate. Aldehyde. Experiment 244. — In a small flask put a few pieces of potassium bichromate, KiOjOt , and pour upon it a few cubic centimetres of moderately concentrated sulphuric acid. To this mixture add slowly a few cubic centimetres of ordinary alcohol. Notice the odor. Soap. Experiment 245. — In an iron pot boil a few ounces of lard with a solution of 40 grams caustic soda in 250 cc. of water for an hour or two. After cooling add a strong solution of sodium chloride. The soap formed will separate and rise to the top of the solution, where it will finally solidify. Dissolve some of the soap thus obtained in water. Hard Water. Experiment 246. — Make some hard water by passing carbon d ioxide through dilute lime-water until the precipitate first formed is dissolved again. Filter. Make a solution of soap by shaking up a few shavings of soap with water. Filter. Add the solution of soap to the hard water. Is a precipitate formed ? Kub a piece of soap between the hands wet with the hard water. Experiment 247. — Make some hard water by shaking a litre or two of water with a little powdered gypsum. Perform with it the same experiments as those first performed with the water con- taining calcium carbonate. CHAPTER XXXVI. OTHER COMPOUNDS OF CARBON. The Carbohydrates. — The carbohydrates form an im- portant group of carbon compounds which include the most abundant substances found in the vegetable kingdom. Besides carbon, they generally contain hydrogen and oxy- gen in the proportions to form water. Hence they are called carbohydrates. The chief compounds included under this head are grape-sugar or glucose, cane-sugar, starch, cellulose, gum, and dextrin. Grape-sugar, Glucose, Dextrose, C.H^O,. — Dextrose occurs very widely distributed in the vegetable kingdom, particularly in sweet fruits. It is found also in honey and, further, in the liver and the blood. Formation of Dextrose. — Dextrose or glucose is formed from several of the carbohydrates by boiling with dilute mineral acids, or by the action of ferments. Its formation from cane-sugar takes place according to this equation, equal quantities of dextrose and levulose being formed: O u H B O u + H 2 = C 6 H ls 6 + C 6 H 12 6 . Cane-sugar. Dextrose. Levulose. Its formation from starch is represented by this equation; C 6 H 10 O 5 + H 3 O = C 6 H la O 6 . Starch. Dextrose. Manufacture of Dextrose or Glucose. — Dextrose is pre pared on the large scale from corn-starch in the United States, and from potato-starch in Germany. The change is usually effected by boiling with dilute sulphuric acid. The acid is afterwards removed by treating with chalk, 667 668 COLLEGE CHEMISTRY. and filtering. [Explain how this removes the acid. J The filtered solutions are evaporated either to a syrupy con- sistency, and sent into the market under the names "glucose," "mixing-syrup," etc.; or to dryness, the solid product being known as "grape-sugar." Properties. — Dextrose crystallizes from concentrated solutions, and as seen in commercial "granulated grape- sugar " looks very much like granulated cane-sugar. It is sweet, but not as sweet as cane-sugar. It is estimated that the sweetness of dextrose is to that of cane-sugar as 3 : 5. Under the influence of yeast it ferments, yielding mainly alcohol and carbon dioxide. Putrid cheese trans- forms it into lactic acid, and then into butyric acid. Levulose, Fruit-sugar, CgH^O,.. — This form of sugar occurs with dextrose in fruits; and is formed by the action of dilute acids or ferments on cane-sugar, which breaks up according to the equation Q ia H M O u + H 2 = C 6 H 12 6 + C 6 H 13 6 . Cane-sugar. Dextrose. Levulose. As cane-sugar is found in unripe fruits, it is probable that the change represented in the equation takes place during the process of ripening. Cane-sugar, C u H 22 O u . — This well-known variety of sugar occurs very widely distributed in nature— in sugar- cane, sorghum, the Java palm, the sugar-maple, beets, madder-root, coffee, walnuts, hazel-nuts, sweet and bitter almonds; in the blossoms of many plants, etc., etc. Sugar-refining. — Sugar is obtained mainly from the sugar-cane and beets. In either case the processes of extraction and refining are largely mechanical. When sugar-cane is used, this is macerated with water to dissolve the sugar. Thus a dark-colored solution is obtained. T^is, is evaporated, and then passed through filters of bone-black by which the color is removed. The clear solution, is then evaporated in open vessels to some extent; and, finally, in large closed vessels called "vacuum-pans," SUGAR, ETC. 669 from which the air is partly exhausted, so that the boiling takes place at a lower temperature than is required under the ordinary pressure of the atmosphere. The mixture of crystals and mother - liquors obtained from the " Tacuum-pans " is freed from the liquid by being brought into the "centrifugals." These are funnel-shaped sieves which are revolved rapidly, the liquid being thus thrown by centrifugal force through the openings of the sieve, while the crystals remain behind and are thus nearly dried. The final drying is effected by placing the crystals in a warm room. Molasses The mother -liquors obtained from the "centrifugals" are further evaporated, and yield lower grades of sugar; and, finally, a syrup is obtained which does not crystallize. This is molasses. Properties of Sugar. — Sugar crystallizes from water in large well-formed prisms. When heated to 210° to 220°, it loses water, and is converted into a substance called- caramel, which is colored more or less brown. When boiled with dilute acids, cane-sugar is split into equal parts of dextrose and levulose. The mixture of the two is called invert-sugar. Yeast gradually transforms cane-sugar into dextrose and levulose, and these then undergo fermenta- tion. Cane-sugar does not ferment. Sugar of Milk, Lactose, C 12 H 22 O n + H 2 0. — This sugar occurs in the milk of all mammals. It is obtained in the manufacture of cheese. Cow's milk consists of water, casein, butter, sugar of milk, and a little inorganij material, in about the following proportions : Water 87 per cent. Casein .• 4 Butter 31 '* Sugar of milk 4f " Mineral matter i Too Cheese is made by adding rennet to milk, which causes the separation of the casein. The sugar of milk remains 6'o COLLEGE CHEMISTRY. in solution, is separated by evaporation, and purified by recrystallization. It has a slightly sweet taste, and is much less soluble in water than cane-sugar. Souring of Milk. — Sugar of milk ferments under certain circumstances, and is transformed mostly into lactic acid. The souring of milk is a result of this fermentation. The lactic acid formed coagulates the casein; hence the thickening. Cellulose, C 6 H ]0 O 6 . — Cellulose forms, as it were, the groundwork of all vegetable tissues. It presents different appearances and different properties, according to the source from which it is obtained; but these differences are due to substances with which the cellulose is mixed; and when they are removed, the cellulose left behind is the same thing, no matter what its source may have been. The coarse wood of trees and the tender shoots of the most delicate plants consist essentially of cellulose. Cotton- wool, hemp, and flax consist almost wholly of cellulose. . Properties. — Cellulose does not crystallize, and is in- soluble in all ordinary solvents. It dissolves in concen- trated sulphuric acid. If the solution is diluted and boiled, the cellulose is converted into dextrin and dextrose. It will thus be seen that rags, paper, and wood, all of which consist largely of cellulose, might be used for the preparation of dextrose or glucose, and consequently of alcohol. Gun-cotton, Pyroxylin, Nitrocellulose. — Cellulose has some of the properties of alcohols ; among them the power to form ethereal salts with acids. Thus, when treated with nitric acid it forms several nitrates, just as glycerin forms the nitrate known as nitroglycerin (which see). The nitrates are explosive, and are used for blasting under the name gun-cotton. Collodion. — A solution of gun-cotton in a mixture of ether and alcohol is known as collodion solution, which is much used in photography. When poured upon any surface, such as glass, the ether and alcohol rapidly evaporate, leaving a thin coating of gun-cotton. STARCH. 671 Celluloid. — Celluloid is an intimate mixture of gun- cotton and camphor. As it is plastic at a slightly elevated temperature, it can easily be moulded into any desired shape. When cooled it hardens. Paper. — Paper in its many forms consists mainly of cellulose. The essential features in the manufacture of paper are, first, the disintegration of the substances used. This is effected partly mechanically and partly by boiling with caustic soda. Then the resulting mass is converted into pulp by means of knives placed on rollers. The pulp, with the necessary quantity of water, is then passed between rollers. Eags of cotton or linen are chiefly used in the manufacture of paper; wood and straw are also used. Starch, C 6 H 10 5 . — Starch is found everywhere in the vegetable kingdom in large quantity, particularly in all kinds of grain, as maize, wheat, etc.; in tubers, as the potato, arrowroot, etc. ; in fruits, as chestnuts, acorns, etc. Manufacture of Starch. — In the United States starch is manufactured mainly from maize; in Europe, from potatoes. The processes made use of are mostly mechan- ical. The maize is first treated with warm water; the softened grain is then ground between stones, a stream of water running constantly into the mill. The thin paste which is carried away is brought upon sieves of silk bolt- • ing-cloth, which are kept in constant motion. The starch passes through with the water as a milky fluid. This is allowed to settle when the water is drawn off. The starch is next treated with water containing a little alkali, the object of which is to dissolve gluten, oil, etc. The mix- ture is now brought into shallow, long wooden runs, where the starch is deposited, the alkaline water running off. Finally, the starch is washed with water, and dried at a low temperature. Properties. — Starch in its usual condition is insoluble in water. If ground with cold water it is partly dissolved. If heated with water the membranes of the cells of which 672 COLLEGE CHEMISTRY. the starch is composed are broken, and the contents form a partial solution. On cooling, it forms a transparent jelly called starch-paste. By dilute acids and ferments starch is converted into dextrin, maltose, and dextrose. Flour.— Wheat flour, which may serve as an example of flour in general, contains water, starch with a little sugar and gum, gluten, and a small quantity of mineral matter. The finest flour contains about 10 per cent of gluten and 70 per cent of starch. Gluten is a substance that resembles in many respects the white of eggs, or egg- albumin. Bread-making. — The chemical changes which take place in bread-making are of special interest. Bread is made by mixing the flour with water and a little yeast. The dough thus prepared is put in a warm place for a time, when it rises. The rising is a result of fermentation caused by the yeast. A part of the starch contained in the flour is converted into sugar, and this is then con- verted into alcohol and carbon dioxide by fermentation. The alcohol passes off for the most part, and the carbon dioxide in striving to escape from the thick gummy dough fills the mass with bubbles of gas, making it light and porous. When the loaf is put into the oven the gases contained in it expand, making it still lighter; then the fermentation is checked by the heat and no further chemical change takes place except on the surface, where the substances are partly decomposed and converted into a dark-colored product, the crust. A FEW COMPOUNDS FROM COAL-TAB. Aromatic Compounds. — The fact that benzene, C 6 H 6 , toluene, 0,H 8 , and other hydrocarbons are obtained from coal-tar has already been mentioned (p. 650). These hydrocarbons are the starting-points for the preparation of a large number of compounds of carbon which are commonly called the "aromatic compounds," as many of them have pleasant aromatic odors. COAL-TAR COMPOUNDS. 673 Nitrobenzene, C 6 H 5 N0 2 .— This substance is formed by treating benzene with nitric acid : C 6 H 6 + HNO, = C 6 H 5 N0 2 + H 2 0. It is a yellow liquid with a pleasant odor like that of the oil of bitter almonds. It is much used under, the name artificial oil of hitter almonds. Aniline, C 6 H 5 NH 2 . — When nitrobenzene is treated with a solution from which hydrogen is given off the oxygen is extracted and replaced by hydrogen: C.H.NO, + 6H = C 6 H 5 NH 2 + 2H 2 0. The product is the substance known as aniline. It is a colorless liquid. When it together with a similar sub- stance, known as toluidine, is treated with mercuric chlo- ride, HgCl 2 , or arsenic acid it is converted into the dye magenta, which is the substance from which most of the aniline dyes are prepared. Aniline Dyes. — Of these a large number are in use. They are all derivatives of rosaniline, of which magenta is a salt. Aniline dyes of great many different colors are made, some of them of great beauty. Phenol, Carbolic Acid, C 6 H 6 0. — This familiar substance is contained in coal-tar, and is extracted from it by treat- ing with caustic soda in which the carbolic acid dissolves. When pure it crystallizes in beautiful colorless rhombic needles. It has a peculiar, penetrating odor, and is poisonous. It is much used as a disinfectant. Oil of Bitter Almonds, \ _, „ .. „,. , , _, . .,, . , I C,H.O. — This substance occurs Benzoic Aldehyde, [ ' 6 in combination with amygdalin, which is found in bitter almonds, laurel-leaves, cherry-kernels, etc. Amygdalin belongs to the class of compounds known as glucosides, which break up into glucose and other substances. Amygdalin itself, under the influence of emulsin, which occurs with it in the plants, breaks up into oil of bitter almonds, hydrocyanic acid, and dextrose : 674 COLLEGE CHEMISTRY. C 20 H a ,NO u + 2H,0 = 0,H 6 + CNH + 20 6 H 12 O s . Amygdalin Oil of Hydrocy- Glucose bitter anic acid almonds It is prepared from bitter almonds, which yield about 1.5 to 2 per cent. It is a liquid which has a pleasant odor. It is made artificially from coal-tar, and is used in the preparation of artificial indigo. Benzoic Acid, C 7 H 6 2 . — Benzoic acid occurs in gum benzoin and in the balsams of Peru and Tolu, and is made artificially from coal-tar by oxidizing toluene,* C,H 8 . Balsams and Odoriferous Resins The balsams of Peru and Tolu are thick fragrant fluids which are obtained from certain trees in South America and elsewhere by cutting the bark. Benzoin is a similar substance. These as well as myrrh, frankincense, and other substances of the kind are used for their odors. The odors are intensified when the substances are heated. They are largely used as incense. Gallic Acid, C,H 6 6 . — Gallic acid occurs in sumach, in Chinese tea, and many other plants. It is formed by boil- ing tannin or tannic acid with suphuric acid. It is pre- pared from gall-nuts by fermentation of the tannin contained in them. It is closely related to tannin or tannic acid. Tannic Acid, Tannin, C u H 10 O 3 . — This substance occurs in gall-nuts, from which it is extracted in large quantities. It is soluble in water. Its solution gives a dark blue-black color with iron salts. Tannin is used extensively in medicine, in dyeing, in the manufacture of leather and of ink. Tanning. — The process of tanning consists in treating hides, from which the hair has been removed, with an infusion of hemlock or oak bark, or of sumach-leaves, in which there is tannic acid. The acid combines with certain parts of the hides, forming insoluble compounds *Tbe name toluene comes from the fact that this hydrocarbon was first obtained from the balsam of Tolu. COAL-TAR COMPOUNDS. 675 which remain in the pores, converting the hides into leather. Indigo. — In several plants which grow in the East and West Indies, in South America, Egypt, and other warm countries, a substance called indican occurs. When this is treated with dilute mineral acids or certain ferments, it breaks up into indigo-blue and a substance resembling glucose. Commerical indigo contains as its principal in- gredient indigo-blue. Indigo- blue is now prepared arti- ficially by the aid of complicated processes. Naphthalene, 10 H 8 . — This hydrocarbon is contained in coal-tar in large quantity. It is a beautiful white crystal- lized substance much used in the preparation of dyes and for protecting woollen fabrics from moth. Anthracene, U H 10 . — Anthracene like naphthalene is obtained from coal-tar. Its chief use is in the preparation of artificial alizarin. Alizarin, O u H 8 4 . — Alizarin is the well-known dye obtained from madder-root. For some years it has been made artificially from anthracene, and the cultivation of madder has been given up. Madder-root was used for dyeing "Turkey-red." Artificial alizarin is almost exclu- sively used for this purpose at present. Glucosides. — Glucosides are substances that occur in nature in the vegetable kingdom. They break down into sugar and other compounds under the influence of fer- ments and dilute acids. Amygdalin has already been mentioned. This breaks down into oil of bitter almonds and dextrose. Indican, which yields indigo and dextrose, is another example. Myronic Acid, another glucoside, is found in the form of the potassium salt in black mustard- seed. When treated with myrosin, which is contained in the aqueous extract of white mustard -seed, potassium myronate is converted into dextrose and oil of mustard. 676 COLLEGE CHEMISTRY. Alkaloids. — These compounds occur in plants, and are frequently those parts of the plants which are most active when taken into the animal body. They are hence some- times called the active principles of plants. Many of these substances arc used in medicine. They all contain nitrogen and in some respects resemble ammonia. Only a few of the more important alkaloids need be mentioned here. Quinine. — This valuable alkaloid is obtained from the outer bark of certain trees which grow in Peru. The bark is known as Peruvian bark. Cocaine is found in cocoa-leaves. Its hydrochloric-acid salt has come into prominence in medicine, owing to the fact that a small quantity of its solution placed upon the eye or the gums or injected beneath the skin causes in- sensibility to pain. Nicotine occurs in tobacco-leaves in combination with malic acid. Morphine, codeine and narcotine are the principal alka- loids found in opium, which is the evaporated sap that flows from incisions in the capsules of the white poppy before they are ripe. INDEX. Acetone, 349. Acetylene, 510, 652, 655 ; se rles, 652. Acid, 156. Acid, acetic, 349, 650, 658 ; an timonic, 333 ; arsenic, 330 arsenic, experiments with, 343 arsenious. 330 ; benzoic, 674 ; boric, 341 ; bromic, 190 ; car- bolic, 673 ; butyric, 659 ; car- bonic, 366 ; chloric, 56, 149 ; chloric, experiments with, 151; chlorous, 149 ; chlorplatinic, 647, 648; chlorsulphuric, 241; chromic, 598; chromic, experi- ments with, 607 ; citric, 661 ; cyanic, 380 ; disilicic, 397; di- sulphuric, 232 ; dithionic, 253; diuranic, 607 ; ferrihydro- cyanic, 630; ferrohydrocyanic, 629; fluosilicic, 392; fluosilicic, preparation of, 400 ; formic, 658 ; fuming sulphuric, 232 ; gallic, 674 ; hydrazoic, 279 ; hydriodic, 192 ; hydrobromic, 188 ; hydrobromic, experi- ments with, 201 ; hydrochloric, 135 ; hydrochloric, experi- ments with, 144 ; hydrochlo- ric, formation of, 144 ; hydro- chloric, preparation of, 144 ; hydrocyanic, 380 ; hydroflu- oric, 197 ; hydrofluoric, ex- periments with, 204 ; hydro- nitric, 279 ; hydrosulphurous, 234 ; hypobromous, 190 ; hy- pochlorous, 148; hyponitrous, 284 ; hypophosphoric, 328 ; hypophosphorous, 329 ; hypo- sulphurous, 234; iodic, 193; iodic, experiments with, 204 ; lactic, 661 ; malic, 661 : man- ganic, 199, 615 ; metaboric, 342 ; metachronous, 595 ; met- antimonic, 321, 333 ; meta- phosphoric, 321, 327 ; metar- senic, 321, 330 ; metastannic, 578, 581 : metavanadic, 339 ; molybdic, 603 ; muriatic, 57 ; myronic, 675 ; nitric, 57, 279 : nitric, experiments with, 293 ; nitric, preparation of, 293 ; nitric, red fuming, 282 ; nitric, reduction of, 295 ; nitrohy- drochloric, 283 ; nitrosylsul- phuric, 225 ; nitrous, 283 ; nitrous, experiments with, 296 ; Nordhausen sulphuric, 232 ; oleic, 662 ; orthoanti- monic, 321, 333 ; orthophos- phoric, 321, 323 ; orthovana- dic, 339 ; osmic, 643 ; oxalic, 660 ; palmitic, 659 ; pentathi- onic, 235; perbromic, 190; perchloric, 149 ; perchloric, preparation of, 152 ; periodic, 173, 195 ; permansranic, 199, 616, 618; persulphuric, 235; phosphoric, 323 ; phosphoric, experiments with, 343 ; phos- phoric, glacial, 328 ; phos- phoric, " insoluble," 505 ; phosphoric, "reverted," 505 ; phosphoric, "soluble," 505; phosphomolybdic, 604 ; phos- phorous, 328 ; propionic, 659 ; prussic, 380 ; pyroantimonic, 321, 333 ; pyroarsenic, 321, 330; pyroligneous, 349, 650; pyrophosphoric, 321, 327; py- rosulphuric. 232; selenic, 242 ; selenious, 242 ; silicic, 394 ; silicic, experiments with, 401 ; silicic, insoluble, 395 ; silicic, 677 678 INDEX. normal, 394 ; stannic, 580 ; stearic, 659 ; sulphantimoni- ous, 336 ; sulpharsenious, 332 ; sulphocyanic, 382 ; sul- phuric, 57, 222 ; sulphuric, experiments with, 246 ; sul- phuric, solid, 232 ; sulphu- rous, 233 ; sulphurous, experi- ments -with, 247, 248 ; sulphy- dric, 211 ; tannic, 674 ; tar- taric, 661 ; telluric, 243 ; tellu- rious, 242 ; tetraboric, 342 ; tetrahydroxylsulphuric, 231; tetrathionic," 235 ; thiosulphu- ric, 234 ; triazoic, 279 ; trithi- onic, 235 ; tungstic, 605 ; va- nadic, 339. Acids, 88, 153 ; basicity of, 161 ; constitution of, 158 ; defini- tion of, 420 ; dibasic, 163 ; fatty, 658 ; monobasic, 162 ; nomenclature of, 166 ; penta- basic, 163 ; polysilicic, 396 ; polystannic, 581 ; tetrabasic, 163 ; tribasic, 163 ; trisilicic, 397. Acids of antimony and araenic, constitution of, 321. Acids of phosphorus, constitu- tion of, 321. Acids of sulphur, constitution of, 223. Acidum phosphoricum glaciate, 328. Affinity, chemical, 9 ; coeffi- cients of, 411. Agate, 393. Air, 253 ; analysis of, 254, 263 ; and life, 259 ; liquid, 261. Alabandite, 614. Alabaster, 503. Albite, 470, 531. Alcohol, 655 ; ethyl, 655 ; me- thyl, 650, 655. Alcohols, 664. Aldehyde, acetic, 657 ; formic, 658. Aldehydes, 657. Algaroth, powder of, 337. Alizarin, 675. Alkalies, 153, 440. Alkali metals, 456. Alkaline earths, 498. Alkaloids, 676. Allanite, 590. Allotropy, 120. Allylene, 652. Alum, 457, 530 ; burnt, 530 shale, 530. Aluminates, 525. Aluminium, 522 ; bronze, 524 ; chloride, 524 ; preparation of, 536 ; hydroxide, 525 ; oxide, 528 ; silicates, 531 : sulDhaie. 528. Alums, 529. Alunite, 530. Amalgamation process for silver. 548. Amalgams, 567. Amethyst, 593. Ammonia, 267, 270; composition of, 274; determination of com- position of, 292 ; experiments with, 290; preparation of, 290. Ammoniacal liquor, 270, 353. Ammonia process for soda, 478. Ammonium alum, 530 ; amal- gam, 277, 568 ; chloride, 483 ; diuranate, 607; hydrosulphide, 485 ; hydroxide, 274 ; nitrate, 485 ; magnesium phosphate, 326, 518; phospho-molybdate, 604 ; salts, 481 ; salts, forma- tion of, 291 ; sulphide, 483. Ammonium sulphide group, 215, 484. Ampere's hypothesis, 102. Anatase, 390, 398. Anhydrides, 194. Anhydrite, 503. Aniline, 673 ; dyes, 673. Anions, 416. Annealing of glass, 508. Anthracene, 650, 675. Antimony, 306 ; and its com- pounds, experiments with, 317 ; blende, 335 ; oxychlo- rides, 336, 344; pentachloride, 313; pentasulphide, 336; pent- oxide, 335 ; salts of, 334 ; sul- phides, experiments with, 344; tetroxide, 335 ; trichloride, 312; trioxide, 333; trisulphide. 335. Antimonyl salts, 334 ; sulphate, 334. Apatite, 300, 494, 504, Aqua regia, 283. Aragonite, 501. INDEX. 679 Arbor Saturni, 584. Argentan, 638. Argentic chloride, 550 ; nitrate, 550. Argentous chloride, 550. Argon, 261. Aromatic compounds, 672. Arrhenius, 422. Arsenic, 305; and its compounds, experiments with, 316 ; disul- phide, 323, 332 ; pentasul- phide, 323, 333 ; pentoxide, 332 ; sulphides, experiments with, 344 ; trichloride, 312 ■ trioxide, 330 ; trioxide, experi- ments with, 344 ; trisulphide, 323, 332. Arsenical pyrites, 634. Arsenides, 305. Arseniuretted hydrogen, 306. Arsenopyrite, 634. Arsine, 306, 316. Atomic theory, 96, 402 ; weights, 17, 19, 105 ; weights and specific heat, 423. Atoms, 102, 403. Aurates, 555. Auric chloride, 557; compounds, 554 ; hydroxide, 554 ; oxide, 554. Aurous chloride, 557 ; com- pounds, 554 ; oxide, 554 ; sul- phide, 557. Avogadro's hypothesis, 102. Azote, 250. Babbitt's metal, 585. Balsams, 674. Banca tin, 578. Barff's process, 626. Barite, 511. Barium, 511 ; carbonate, 513 ; chloride, 511 ; chromate, 602, dioxide, 512 ; hydroxide, 511 ; oxide, 511 ; peroxide, 512 ; phosphates, " 513 ; sulphate, 513 ; sulphide, 513. Base 1 57 Bases, 88, 153 ; acidity of, 163 ; constitution of, 158 ; definition of, 420 ; diacid, 163 ; monacid, 163 ; nomenclature of, 167 ; triacid, 163. Basic-lining process, 625. Battery, storage, 587. Bauxite', 522. Bell-metal, 540. Bengal fire, 511. Benzene, 356, 652, 672 ; series, 652. Benzine, 651. Benzoin, 674. Berthollet, 126. Beryl, 515. Beryllium, see Glucinum. Bessemer process, 625. Bismuth, 309 ; basic nitrates of, 337, 344 ; dichloride, 313 ; di- oxide, 338 ; experiment with, 317; nitrate, 310; oxy chloride, 338; pentoxide, 338; sub- nitrate, 338 ; sulphate, 310 ; trichloride, 313 ; trioxide, 337; trisulphide, 338. Bismuthyl salts, 337. Blast furnace, 622. , Black ash, 477. Black lead, 348. Bleaching by chlorine, 130 ; by sulphur dioxide, 237 ; powder, .498. Block tin, 578. Blow-pipe, 376 ; experiments with, 384. Bluestone, 543. Blue vitriol, 543, 544. Bog iron ore, 631 Bone-black, 351 ; filters, 352, 357 ; oil, 650. Bones, distillation of, 650. Boracite, 341, 519. Borax, 339,' 342, 480. Boron, 339 ; experiments with, 344 ; nitride, 343 ; phosphate, 342 ; salts of, 342 ; trichloride, 340. Boryl potassium tartrate, 342. Boussingault, 255. Boyle's law, 44. Brass, 540, 561. Braunite, 610, 612. Bread-making, 672. Bricks, 534. Brimstone, crude, 207 ; roll, 207. Britannia metal, 307, 579. Bromides, experiments with, • 452. Bromine, 187 ; preparation of, 201 ; water, 188. 68o INDEX. Bronze, 540. Brookite, 390, 398. Brown iron ore, 622. Bunsen, 256, 257, 469. Bunsen burner, 378. Burning, 36. Butane, 652. Butter, 663 ; of antimony, 313. Butylene, 652. Butyrum Antimonii, 313. Cadmium, 564 ; carbonate, 565 ; chloride, 565 ; sulphate, 565 ; sulphide, 565. Caesium, 469, 470. Calcium, 494 ; carbide, 510 carbonate, 501 ; chloride, 495 chloride, preparation of, 520 fluoride, 496 ; hydroxide, 496 nitride, 509 ; oxide, 496; phos- phates, 504 ; silicate, 506 ; sul- phate, 503 ; sulphide, 509. Calc-spar, 501. Calomel, 568. Calorie, 40. Calorimeter, 39. Candle, standard, 373. Caue-sugar, 668. Caramel, 669. Carbohydrates, 667. Carbon, 345 ; amorphous, 348 ; chemical conduct of, 355 ; di- oxide, 359 ; dioxide, experi- ments -with, 369 ; dioxide, re- lations to life, 364;disulphide, 381 ; experiments on reducing power of, 358; monoxide, 367; experiments with, 371 ; sili- cide, 398. Carbonates, 366, 449 ; acid, 366 ; basic, 367 ; experiments, 455. Carbonyl chloride, 369. Carbon silicide, 398. Carborundum, 398. Carnallite, 457, 459, 515. Carnelian, 393. Caro's reagent, 235. Cassiterite, 577. Cast iron, 623; gray, 624; white, 623. Cations, 416. Caustic soda, 473. Cavendish, 54. Celestite, 510, 511. Celluloid, 671. Cellulose, 667, 670. Cements, 509 ; hydraulic, 509. Cerite, 590. Cerium, 390, 590, 591. Cerussite, 583. Chalcedony, 393. Chalcocite, 538; 545. Chalk, 494, 501, 502. Chance process for recovery of sulphur, 478. Charcoal, 348 ; absorption of gases by, 351, 357 ; animal, 351 ; filters, 352, 357 ; kiln, 349. Charring process, 348. Chemical action, 7. Chemical analysis, 648. Chemical change, 2 ; caused by heat, 26. Chemical reactions, cause of, 409; kinds of, 405 ; ideal, 409. Chemistry, 3^ ; organic, 346, 649. Chili saltpetre, 191, 279, 280, 473. Chloral, 675 ; hydrate, 657. Chlorates, 447. Chloraurates, 555. Chlorides, 130, 432; double, 437; experiments with, 452; general properties of, 436. Chlorination process, 556. Chlorine, 126, 142; dioxide, 150; experiments with, 142 ; prepa- ration of, 127; heptoxide, 151; hydrate, 133, 143; liquid, 133; monoxide, 150. Chloro-acids, 437. Chloroform, 654. Chlorostannates, 580. Chloroform, 654. Chlorophyll, 622. Chlorplatinates, 647. Choke-damp, 364. - • Chromates, 598 ; experiments with, 607. Chrome alums, 598. Chrome yellow, 602. Chromic chloride, 596 ; com- pounds, 594 ; hydroxide, 596 ; iron, 595 ; oxide, 597 ; sul- phate, 597. Chromite, 527, 595. Chromium, 244, 594; trioxide, 600. Chromous compounds, 594. INDEX. 681 Chrysoberyl, 526. Cinnabar, 567, 572. Clay, 388, 522, 532. Cleveite, 536. Coal, 352 ; anthracite, 352 ; bitu- mina»js, 352; gas, 372, 383; dis- tillation of, 649 ; tar, 353, 649. Cobalt, 635 ; sulphide, 636. Cobaltic hydroxide, 636 ; oxide, 636. Cobaltite, 636, 637. Cobaltous chloride, 636 ; cya- nide, 637 ; hydroxide, 636. Cocaine, 676. Codeine, 676. Coke, 350. Collodion, 670. Columbite, 339. Columbium, 339. Combination, 22 ; direct, 405. Combining weights, 15. Combustion, 37 ; experiments on, 52. Compound, chemical, 28. Condy's liquid, 618. Constitution, 108, 404. Copper, 536 ; acetate, 658 ; basic carbonate, 545 ; metallurgy of, 538 ; native, 538 ; plating, 545 ; ruby, 538, 542. Copperas, 633. Coral, 494. Corrosive sublimate, 568. Corundum, 528. Cotunnite, 585. Crafts, 524. Crocoisite, 583, 595. Cryolite, 470, 522. Cupric arsenite, 544 ; carbonates, 545 ; chloride, 541 ; hydrox- ide, 542 ; nitrate, 544 ; oxide, 542 ; sulphate, 543 ; sulphide, 545. Cuprite, 542. Cuprous chloride, 541 ; chloride, preparation of, 558 ; hydrox- ide, 542; oxide, 541; sulphide, 545. Cyanides, 379. Cyanogen, 379 ; experiments with, 386. Dalton, 96. Dalton's law, 44 Datholite, 341. Davy, 126. Deacon's process, 127. Decay, 458. Decomposition, 23 ; direct, 407 ; double, -23, 408. Decrepitation, 472. De la Bastie glass, 508. Deliquescence, 84, 95. Developers in photography, 552. Deville, 524. Dextrin, 667. Dextrose, 667. Dialyser, 395. Dialysis, 395. Diamond, 347. Diaspore, 525. Didymium, 339, 591. Diffusion, 58, 260 ; experiments on, 68. Dimorphism, 209. Dissociation, 87, 311, 414 ; elec- trolytic, 417; of dissolved sub- stances, 417. Distillation, 95 ; destructive, 349 ; dry, 349. Dolomite, 359, 515. Double chloride, 437. Drummond light, 78. Dulong, 426. Dumas, 255 ; his study of com- bustion of hydrogen, 72 ; method for determining the specific gravity of vapors, 112. Dynamite, 663. Earthenware, 534. Efflorescence, 84, 95, 474. Bka-aluminium, 574. Eka-boron, 535. Eka- silicon, 577. Electrolysis, 416. Electrolytes, 416. Electrolytic dissociation, 417. Electrolytic process for chlorine, 129. Electro-negative ions, 417. Electro-positive ions, 417. Electrotypes, 546. Elements, 6, 17 ; acid-forming, 141 ; base-forming, 158 ; re- placing power. 111 ; symbols of, 19. Emerald, 51 5. Emery, 528. Emulsin, 673. 682 INDEX. Energy, chemical, 41; conserva- tion of, 5; stored up in plants, 365. Enstatite, 397. Epsom salt, 515, 517. Equations, chemical, 22. Equilibrium, 412. Erbium, 520. Etching on glass, 198, 202. Ethane, 652. Ether, 661. Ethereal salts, 662. Ethers, 661. Ethiops martialis, 572. Ethylene, 654 ; series, 652. Eudiometer, 73. Eudiometric method, 73 ; ex- periments, 81. Faraday's law, 416. Fats, 662. Feldspar, 451, 457, 522, 531. Ferment, nitrifying, 280. Fermentation, 650, 656 ; acetic acid, 656 ; alcoholic, 656 ; lac- tic acid, 656. Ferric chloride, 627 ; chloride, preparation of, 640 ; ferro- cyanide, 630 ; hydroxide, 630 ; soluble, 631 ; oxide, 632 ; sul- phate, 633 ; sulphate, prepara- tion of, 640. Ferrous ammonium sulphate, preparation of, 640 ; carbon- ate, 632 ; chloride, 627, ; ferric oxide, 631 ; ferricyanide, 630 ; hydroxide, 630 ; phosphate, 633 ; sulphate, 633 ; sulphate, preparation of, 640 ; sulphide, 632. Fire-damp, 653. Flame, 373 ; oxidizing, 375 ; re- actions, 487 ; reducing, 375. Flames, 373 ; causes of lumin- osity of, 377 ; structure of, 375 Flint, 393. Mores zinci, 564. Flour, 672. Fluorine, 196. Fluor-spar, 494, 496. Flux, 197, 496, 502. Fool's gold, 634. Formulas, constitutional, 109 ; molecular, 107, 403. Frankincense, 674. Franklinite, 560, 631. Fulminating mercury, 572 . Fumaroles, 341. Fusel oil, 656. Gadolinite, 590. Gahnite, 526, 560. Galena, 583. Galenite, 583. Galvanized iron, 561. Gallium, 574. Garnet, 506 Gas, normal, 104. Gases, kindling temperature of, 374 ; measurement of volume of, 43. Gasoline, 651. Gay Lussac, 126 ; law of, 44 ; tower, 227. Germanium, 577 ; chloride, 577. German silver, 541, 638. Gersdorffite, 638. Glass, 506. Glauber, 135, 474. Glauber's salt, 474. Glover tower, 227. Glucinum, 515. Glucose, 667. Glucosides, 675. Gluten, 672. Glycerin, 657. Gold, 554 ; alloys, 557 ; amal- gam, 557 ; metallurgy of, 555 ; sulphide, 557. Goldschmidt's welding process, 523. Graham, 61. Granite, 531. Grape-sugar, 667. Graphite, 347. Greenockite, 565. Green vitriol, 544, 638. Gul Iberg, 411. Gum, 667. Gun-cotton, 670. Gun-metal, 540. Gunpowder, 463. Gypsum, 448, 494, 503. Halogens, 186. Hardnsss, permanent, 502, 504 • temporary, 502. Hard water, 502 ; experiments with, 666. INDEX. 683 Hauerite, 614. Hausmannite, 610, 612. Heat, of combustion, 39 ; of de- composition, 40; of neutraliza- tion, 418 ; specific, 4aw of, 425. Heavy spar, 448, 511. Helium, 536. Hematite, 622, 632. Heptane, 652. Hexane, 652. Homologous series, 652. Homology, 652. Hornblende, 515. Hydrargilite, 525. Hydrates, 87, 167, 169. Hydraulic cement, 509 ; mining, 555. Hydrazine, 278. Hydrazoic acid, 279. Hydrocarbons, 356, 358, 650; formation of, 649. Hydrochloric acid hydrate, 138. Hydrogen, 54 ; amount evolved when a known weight of a metal is dissolved in an acid, 65 burning of, 72; chemical prop erties of, 59, 69 ; dioxide, 122 dioxide, experiments with 124 ; peroxide, 122 ; persul phide, 217; preparation of, 54. 62 ; sulphide, 211 ; sulphide, experiments with, 221 ; sul- phide group, 214 ; sulphide in chemical analysis, 214. Hydrogenium, 61. Hydrogen- valence, 182. Hydrosulphides, 216, 445. Hydroxides, 440 ; experiments with, 453 ; formation from oxides, 239. Hydroxyl, 664. Hydroxylamine, 278. Hypochlorites, 447. Hyposulphite of soda, 475. Iceland spar, 501. Ice-machine, 271. Illuminating gas, 372. Illumination, 372. Incense, 674. Indican, 675. Indigo, 675 Indium, 575. Infusorial earth, 394 Ink, sympathetic, 636. Invert- sugar, 669. Iodic anhydride, 194. Iodides, experiments with, 452. Iodine, 190; bromide, 196 ; chlo- ride, 196 ; experiments with, 202 ; pentoxide, 194 ; trichlo- ride, 196. Iodoform, 654. Ions, 90, 416, 417. Iridium, 644. Iron, 621 ; disulphide, 634 ; gal vanized, 561; metallurgy of, 622 ; ores, 622 ; pyrites, 634. Kainite, 466, 515. Kaolin, 522, 532. Kelp, 190. Kieserite, 515, 517. Kindling temperature, 37 ; of gases, 374. King's yellow, 332. Kirchhoff, 469. Kruegel, 262. Krypton, 262. Lactose, 669. Ladenburg, 262. Lamp-black, 350. Lanthanum, 536, 590, 591. Lapis lazuli, 533. Laughing-gas, 285. Lavoisier, 32, 250, 254. Law of definite proportions, 12, 402; of Dulong and Petit, 425 ; of indestructibility of matter, 4, 402 ; of multiple proportions, 13, 402 ; of speci- fic heats, 425. Lead, 582 ; acetate, 658; carbon- ate, 588 ; chloride, 585, 592 ; chromate, 602 ; hydroxide, 586 ; iodide, 585, 592 ; metal- lurgy of, 583; molybdate, 603 ; nitrate, 588 ; oxide, 585 ; pen- cils, 348 ; peroxide, 586, 592 ; sesquioxide, 592 ; suboxide, 585 ; sugar of, 658 ; sulphate, 589 ; sulphide, 588 ; tetrachlo- ride, 585 ; tree, 584. Le Blanc process for soda, 476. Lepidolite, 469, 481. Levulose, 668. Lewes, 378. 68 4 INDEX. Lignite, 352. Lime, 496 ; air-slaked, 496 ; chloride of, 498 ; slaked, 496. Lime-kiln, 496. Lime-light, 78. Limestone, 494, 501. Lime water, 497. Linking of atoms, 404. Linnaeite, 637. Liquor ferri ehlorati, 627 ; ferri sesquichlorati, 628. Litharge, 586. Lithium, 481 ; carbonate, 481 ; chloride, 481 ; phosphate, 481. Loadstone, 631. Lunar caustic, 553. Magenta, 673. Magnalium, 534. Magnesia, 517 ; alba, 518 ; wsta, 517. Magnesio-ferrite, 527. Magnesite, 515. Magnesium, 515 ; borates, 519 carbonate, 518 ; chloride, 516 chloride, preparation of, 520 hydroxide, 517 ; oxide, 517 phosphates, 518; silicates, 519 silicide, 519 ; sulphate, 517. Magnetite, 527, 622, 631. Malachite, 538, 545. Malaria, 261. Manganates, 615. Manganese, 199, 609 ; black oxide of, 199, 610, 612 ; di- oxide, 199, 610, 612; disul- phide, 614 ; heptoxide, 201, 618 ; salts, preparation of, 619 ; tetrachloride, 611 ; tetra- fluoride, 611 ; trichloride, 611. Manganic oxide, 612 ; sulphate, 615. Manganite, 610, 612. Manganous chloride, 610, 619 ; hydroxide, 612 ; manganic oxide, 612 ; sulphate, 614 ; sulphide, 614. Marble, 494, 501. Marcasite, 634. Marl, 501, 532. Marsh- gas, 356, 653. Marsh's test for arsenic, 308, 317. Mass action, 411. Mass, influence of, 410. Matches, safety, 303. Matter, constitution of, 5 ; law of indestructibility of, 4. Meerschaum, 515. MendeleefE, 173. Mercuric chloride, 569 ; com- pounds, 566 ; .cyanide, 572 ; diammonium chloride, 573 ; iodide, 570 ; nitrate, 573 ; ni- trate, preparation of, 575 ; ox- ide, 571 ; sulphide, 572. Mereurius solubilis Hahne- manni, 573. Mercurous ammonium chloride, 573 ; chloramide, 573 ; chlo- ride, 568 ; compounds, 566 ; iodide, 570 ; nitrate, 573 ; ni- trate, preparation of, 575. Mercury, 566 ; metallurgy of, 567. Meta-arsenates, 330. Metallic properties, 428. Metallurgy, 430. Metals, 158, 428 ; properties, 431. Metaphosphates, 326. Metastannates, 581. Metathesis, 23, 408. Meteorites, 621, 637. Methane, 356, 653. Methyl alcohol, 349. Meyer, Lothar, 173. Meyer, Victor, method for de- termining the specific gravity of vapors, 113. Mica, 506, 522, 531. Microcosmic salt, 485. Milk, souring of, 670. Minerals, 430. Miner's safety lamp, 374. Minium, 587. Mixtures, mechanical, 10, 28. Moissan, 347. Molasses, 669. Molecular weights, 103. Molecules, 102, 403. Molybdates, 603. Molybdenite, 602. Molybdenum, 244, 602. Morley's determination of the ratio between the atomic weights of hydrogen and oxy- gen, 73. Morphine, 076. fNDEX. 685 Mortar, 508. Mosaic gold, 582. Murium, 126. Myrrh, 674. Naphtha, 651. • Naphthalene, 650, 675. Narcotine, 676. Nascent state, 121. Neodymium 339, 591. Neon, 262. Neutralization, 153 ; experiments on, 170. Newton's metal, 310. Niccolite, 638. Nickel, 637 ; alloys, 638 ; car- bonyl, 639; cyanide, 638; plating, 638. Nickelic hydroxide, 638. Nickelous hydroxide, 638 ; ox- ide, 638. Nicotine, 676. Nilson, 524. Niobium, 339. Nitrates, 446 ; formation of, 268 ; formation of, in the soil, 268, 280, 462. Nitric oxide, 285 ; experiments with, 296. Nitrification, 280, 462. Nitrobenzene, 673. Nitrocellulose, 670. Nitrogen, 250 ; boride, 343 ; chloride, 289 ; experiments with, 262 ; iodide, 290 ; pent- oxide, 289 ; peroxide, 287 ; peroxide, ' experiments with, 297 ; preparation of, 251, 262; relations to life, 259 ; trioxide, 287 ; trioxide, experiments with, 297. Nitroglycerin, 663. Nitrosyl chloride, 283. Nitrous oxide, 284, experiments with, 296. Non-metals, 428. Octane, 652. Oil of bitter almonds, 673 ; arti- ficial, 673. Olefiant gas, 356, 654. Oleomarearin, 663. Olivine, 519. Opal, 393. Opium, 676. Ores, 430. Orpiment, 332. Orthoclase, 398. Osmium, 643 ; peroxide, 643. Osmotic pressure, 422. Oxidation, slow, 38. Oxides, 42, 439 ; acidic, 194 ; basic, 194. Oxygen, 32 ; amount liberated from a known weight of po- tassium chlorate, 47 ; burning of, 60 ; chemical properties of, 35, 50 ; physical properties of, 35, 50 ; preparation of, 32, 42 ; valence, 182. Oxyhydrogen light, 78, 82; blow- pipe, 77 ; blow-pipe, experi- ments with, 81. Ozone, 116 ; in the air, 119. Palladium, 644 ; hydrogen, 644. Paper, 671. Paraffin, 651. Paraldehyde, 657. Paris green, 544. Passive state of iron, 627. Pattinson's method of separating silver from lead, 548. Peat, 352. Pentane, 652. Periodates, 195. Periodic law, 173, 405. Permanent white 513. Permanganates, 616. Petalite, 481. Petit, 425. Petroleum, 649, 651. Pettersson, 524. Pfeffer, 422. Phenol, 673. Philosopher's wool, 563. Phosgene, 369. Phosphates, 450; acid, 325 neutral, 325 ; normal, 325 primary, 325 ; secondary, 325 tertiary, 325. Phosphine, 304 ; experiments with, 315 ; liquid, 305. Phosphoric anhydride, 329. Phosphorite, 300, 494, 504 Phosphorous anhydride, 329. Phosphorus, 300 ; experiments with, 314 ; oxychloride, 330 ; pentachloride, 311 ; pentachlo- 686 INDEX. ride, preparation of , 819; pent- oxide, 329 ; red, 303 ; trichlo- ride, 310 ; trichloride, prep- aration of, 317 ; trioxide, 329. Phosphorus, constitution of acids of, 321. Phosphuretted hydrogen, 304. Photography, 552. Photometer, 373. Physical change, 2. Physics, 3. Pig-iron, 623. Pinchbeck, 540. Pink salt, 580. Pitch blende, 605. Placer mining, 555. Plaster of Paris, 503. Platinic chloride, 647 ; sulphide, 648. Platinous chloride, 647; sulphide, 648. Platinum, 645 ; alloys, 647 ; black, 646 ; metals, 642 ; spongy, 645. Plumbago, 347. Plumbates, 587. Plumbites, 586. Pollux, 469. Porcelain, 532, 533. Potash, 457, 467. Potassium, 457 ; alum, 530 -, bi- chromate, 599, 607 ; bromide. 459; carbonate, 467; carbonate, acid, 468 ; carbonate, extrac tion from wood-ashes, 490 chlorate, 464 ; chlorate, prep aration of, 151 ; chloride, 459 chromate, 599, 607 ; cyanide, 465 ; f erricyanide, 630 ; f erro cyanide, 379, 629 ; fluoger manate, 577 ; fluoride, 459 fluosilicate, 487; hydride, 458 hydrosulphide, 461 ; hydrox ide, 460 ; hypochlorite, 147 iodide, 459 ; iodide, prepara tion of, 489 ; manganate, 615 manganate, preparation of. 619 ; nitrate, 462 ; nitrite 463 ; oxide, 461 ; perchlorate, 465 ; perchlorate, preparation of, 152 ; permanganate, 200. 617 ; permanganate, prepara tion of, 619 ; peroxide, 461 phosphates, 468 ; silicate, 469 sulphate, 466 ; sulphate, acid, 466 ; sulphite, 467 ; sulphite, acid, 467 ; sulphocyanate, 466. Praseodymium, 339, 591. Preparing salt, 581. Priestley, 32, 254. Printing-ink, 350. Propane, 652. Propylene, 652. Prussian blue, 630. Puddling, 624. Purple of Cassius, 55'?. Pyrargyrite, 553. Pyrite, 622, 634. Pyroarsenates, 330. Pyrolusite, 199, 610, 612. Pyrophosphates, 327. Pyrosiderite, 631. Pyroxylin, 670. Pyrrhotite, 622. Quartation, -556. Quartz, 388. Quartzite, 388, 393 Quicklime, 496. Quinine, 676. Radicals, 664. Ramsay, 261, 262. Rayleigh, 261. Raoult's methods for determin- ing molecular weights, 420, 422. Reaction, chemical, 409. Realgar, 332. Red fire, 511. Red lead, 587. Red phosphorus, 303. Red prussiate of potash, 638. Reducing agent, 60. Reduction, 60, 70 ; by carbon, 356. Residues, 664. Respiration. 363. Reversion of phosphates, 505. Rhodium, 644. Rhodocroisite, 610. Rinmann's green, 564. Roasting, 445 Rock-crystal, 393. Rose's metal, 310. Rouge, 632. Rubidium, 469, 470. INDEX. 687 Ruby, 528. Ruby copper, 538, 542. Rupert's drops, 508. Ruthenium, 643. Rutile, 390, 398. Safety lamp, 374. Safety matches, 303, 465. Sal ammoniac, 270, 483. Salt, common, 471. Saltpetre, 457, 462 ; plantations, 462. Salts, 157 ; acid, 162, 164 ; basic, 163, 164 ; constitution of, 160 ; decomposition by bases, 441 ; formation of, 434 ; neu- tral, 162 ; nomenclature of, 168 ; normal, 162, 164. Samarium, 536. Samarskite, 536. Sand, 388, 393. Saponification, 662. Sapphire, 528. Scandium, 535. Scheele, 32, 126, 254. Scheele's green, 544. Scheelite, 604. Schlippe's salt, 336. Schweinfurt green, 544. Selenium, 218 ; dioxide, 242. Serpentine, 515. Siderite, 622, 632. Siemens-Martin furnace, 625. Silica, 388. Silicates, 388, 451. Silicides, 398. Silicon, 388 ; dioxide, 388, 393 ; hexachloride, 391 ; hydride, 390 ; magnesium, 519 ; prepa- ration of, 399 ; tetrachloride, 391 ; tetrafluoride, 392 ; tetra- fluoride, preparation of, 400. Silver, 546 ; allotropic forms of, 549 ; alloys of, 550 ; ainal gam, 550 ; bromide, 550 chloride, 550 ; cyanide, 553 iodide, 551 ; metallurgy of. 547 ; nitrate, 553 ; oxide, 552 peroxide, 553 ; suboxide, 553 tree, preparation of, 558 triazoate, 552. Slaking, 497. Smalt, 637. Smaltite, 636. Smithsonite, 560, 564 Soap, action on hard waters, 659 ; preparation, 666. Soaps, 659. Soapstone, 515. Soda, 475 ; calcined, purified, 477 ; crude, 477 ; crystallized, 477. Soda-water, 362. Sodium, 470 ; alum, 531 ; amal- gam, 567 ; ammonium phos- phate, 485 ; bichromate, 602 ; borate, 480 ; bromide, 473 ; carbonate, 475 ; carbonate, acid or primary, 479 ; carbo- nate, preparation of by am- monia process, 478 ; chloride, 471; chromate, 602; diuranate, 607 ; fluoride, .473; hydride, 471 ; hydroxide/ 473 ; iodide, 473 ; oxide, 473 ; nitrate, 473 ; peroxide, 473 ; phosphates, 479 ; silicate, 480 ; stannate, 580 ; sulphantimonate, 473 ; sulphate, 473 ; sulphate, ex- periment on supersaturated solution of, 490 ; thiosulphate, 475. Solder, soft, 579. Solution, 88 ; as an aid to chem- ical action, 90. Solvay process for soda, 478. Spathic iron, 632. Specific heat and atomicweights, 423. Specific gravity of vapors, de- termination of, 112. Spectroscope, 487. Spectrum, continuous, 488; line, 488. Spelter, 560. Sphalerite, 560. Spiegel iron, 624. Spinels, 526. Spirits of wine, 655. Spiritus fumans IAbcmi, 580. Spitting of silver, 549. Spodumene, 481. Stalactites, 502. Stalagmites, 502. Stannic chloride, 580 ; chloride preparation of , 592 ; hydrox- ide, 580 ; oxide, 581 ; sul- phide, 581. Stannite, 578. Stannous chloride, 579; chloride. 688 INDEX. preparation of. 592 ; sulphide, 581. Starch, 667, 671. Stearin, 659. Steel, 625. Stibine, 307, 317. Stibnite, 306, 335. Stolzite, 605. Storage battery, 587. Strass, 507. Stromeyerite, 553. Strontianite, 510. Strontium, 510 ; chloride, 510 ; hydroxide, 510 ; nitrate, 510 ; oxide, 510 ; sulphate, 511. Sublimation, 483. Substitution, 55. Sugar of lead, 658; of milk, 669; refining, 668. Sulphantimonates, 336. -Sulphates, 448 ; experiments with, 454. Sulphides, 210, 443. Sulphites, 449 ; acid, 233. Sulpho-salts, 446. Sulphostannates, 582. Sulphur, 206 ; acid chlorides of, 240; auratum,336; dichloride, 218 ; dioxide, 235 ; dioxide, ex- periments with, 247 ; flowers of, 207; heptoxide, 235 ; hexa- fluoride, 218 ; insoluble, 210 ; monochloride, 218 ; sesquiox- ide, 235 ; soluble, 210 ; stick, 207 ; tetrachloride, 218 ; triox- ide, 238; trioxide, experiments with, 248 ; waters, 212. Sulphur, constitution of acids of, 223. Sulphuretted hydrogen, 211. Sulphuryl chloride, 240. Sulphuryl hydroxyl chloride, 241. Sulphydrates, 217. Superphosphate of lime, 505. Sylvite, 457, 459. Symbols of compounds, 20 ; of elements, 19. Sympathetic inks, 636. Tachydrite, 495. Talc, 515. Tannin, 674. Tanning, 674. Tantalite, 339. Tantalum, 339. Tartar, cream of, 661 ; crude, 457 ; emetic, 334. Tellurium, 219; dioxide, 244; monoxide, 244 ; trioxide, 244. Tempering of glass, 508; of steel, 625. Thallium, 575. Thenard, 126. Theory, use and value of a, 99. Thomas- Gilchrist process, 625. Thorite, 390, 397. Thorium, 390. Tin, 577 ; amalgam, 579 ; metal- lurgy of, 578 ; salt, 579 ; stone, 577, 581. Titanium, 390 ; dioxide, 395. Toluene, 652, 672. Toluidine, 673. Tridymite, 393. Troost, 524. Tungstates, 605. Tungsten, 244, 604. Turkey red, 675. Turnbull's blue, 630. Tuyeres, 622. Type metal, 307. Ultramarine, 533. Uranates, 607. Uraninite, 606. Uranium, 244, 605 ; oxides, 606 ; yellow, 607. Uranous salts, 606. Uranyl nitrate, 606 ; sulphate, 607 ; salts, 606. Valence, 109, 404 ; variations of, 404. Vanadium, 338. Van't Hoff, 423. Vein-mining, 555. Ventilation, 260. Verdigris, 658. Vinegar, 658 ; mother of, 656, 658. Vitriol, oil of, 57, 223. Vitriols, 544. Vivianite, 633. Waage, 411. "Water, 83; as a solvent, 88; con- tamination of, by sewage, 92 ; INDEX. 689 determination of composition of, 79; in organic substances, 83, 94 ; of crystallization, 83* 94 ; proofs of composition of, 84 ; purification of, 93, 95 ; synthesis of, 85. Water-gas, 368. Water-glass, 469, 481. Waters, chalybeate, 92 ; effer- vescent, 93 ; natural, 91 ; sul- phur, 92. Weights, atomic, 17, 100 ; com- bining, 15, 100. Weldon's process, 128, 613. Welsbach, Auer von, 591. White arsenic, 330. White lead, 588. White precipitate, 573 ; fusible, 573 ; infusible, 573. White vitriol, 544, 564. Winkler, 576. , Witherite, 511. Wolframite, 604. Wollastonite, 397, 506. Wood, distillation of, 650. Wood's metal, 310. Wood spirit, 349, 650, 655. Work, chemical, 41. Wrought iron, 624, 625. Wulfenite, 583, 603. Xenon, 262. Xylene, 652. Yellow prussiate of potash, 379, 628. Ytterbium, 536. Yttrium, 536. Zeolites, 506. Zinc, 560 ; alloys, 561 ; amal- gam, 561 ; blende, 563 ; car- bonate, 564 ; chloride, 562 ; dust, 560 ; experiment on burning of, 575 ; hydroxide, 562 ; metallurgy of, 560 ; ox- ide, 563 ; solution of, in so- dium hydroxide, 575 ; sul- phate, 564 ; sulphide, 563 ; white, 563. Zircon, 390, 397. Zirconium, 390.