DESCRIPTIVE CHEMISTRY BY LYMAN C. NEWELL, PH.D. (JOHNS HOPKINS) INSTRUCTOR IN CHEMISTRY, STATE NORMAL SCHOOL, LOWELL, MASS. AUTHOR OF "EXPERIMENTAL CHEMISTRY" BOSTON, U.S.A. D. C. HEATH & CO., PUBLISHERS 1903 COPYRIGHT, 1903, BY LYMAN C. NEWELL. ANTOIINE LAURENT LAVOISIER 1743-1794 THE CELEBRATED FRENCH CHEMIST WHO LAID THE FOUNDATIONS OF CHEMISTRY PREFACE. THIS book is intended for teachers who wish to emphasize the facts, laws, theories, and applications of chemistry. It is divided into two parts. Part I contains the text, together with exercises and problems. Part II contains the experiments. The text has been selected and arranged with special refer- ence to the needs of teachers as well as to the capacity of students. The experiments have been prepared to meet the needs of those schools in which the laboratory facilities are limited or the time for chemistry is short. The point of view differs from that in the author's " Experimental Chem- istry," but the spirit is the same. The two books are companion volumes, though of course they' can be used independently. The cordial reception given the " Experimental Chemistry " shows that many teachers are empha- sizing the experimental side of chemistry. These teachers will find Part I of the " Descriptive Chemistry " a serviceable companion book both in the laboratory and class room. It has been bound as a separate volume to meet such a use. Solutions of problems, answers to some of the exercises, and references to the literature have been put in a separate Teacher's Handbook. The manuscript has been read by Dr. William B. Schober, Lehigh Uni- versity, Bethlehem, Pennsylvania; Mr. Franklin T. Kurt, Chauncey Hall School, Boston, Massachusetts; and Mr. George M. Turner, Masten Park High School, Buffalo, New York. The chapters on theory were also read by Dr. Alexander Smith of the University of Chicago, and the chapters on carbon by Dr. James F. Norris of the Massachusetts Institute of Technology. The proof has been read by Dr. E. H. Kraus, High School, Syracuse, New York; Professor E. S. Babcock, Alfred University, Alfred, New York; and Mr. E. R. Whitney, High School, Binghamton, New York. The author is grateful to these teachers for their criticism, but he assumes all responsibility for any errors which may be detected. L. C. N. LOWELL, MASS., 239204 iii ,... CONTENTS. PART I. , f } CHAPTER PAGE -I. PHYSICAL AND CHEMICAL CHANGES CHEMICAL ACTION, ~r CHEMICAL ENERGY ELEMENTS COMPOUNDS . '. i II. OXYGEN LAWS OF CHARLES AND BOYLE OZONE . . 1 1 III. HYDROGEN . . . .... . . . .23 IV. GENERAL PROPERTIES OF WATER 31 V. COMPOSITION OF WATER HYDROGEN DIOXIDE ... 50 fVI. THE ATMOSPHERE NITROGEN 61 VII. LAW AND THEORY LAWS OF DEFINITE AND MULTI-J^ PROPORTIONS ATOMIC THEORY ATOMS AND MOLE- CULES SYMBOLS AND FORMULAS EQUATIONS . . 75 VIII. ACIDS, BASES, AND SALTS 87 IX. EQUIVALENTS ATOMIC AND MOLECULAR WEIGHTS CHEMI- CAL CALCULATIONS QUANTITATIVE SIGNIFICANCE OF EQUATIONS .... -^^^^^- 100 X. LIGHT, HEAT, ELECTRICITY, AND CHEMICAJMPN . . in XI. CHLORINE AND HYDROCHLORIC ACID . . . . .133 XII. AMMONIA NITRIC ACID AND NITRATES AQUA REGIA ^p OXIDES OF NITROGEN ^ 14*1 XIII. GASES GAY-LUSSAC'S LAW AVOGADRO'S HYPOTHESIS VAPOR DENSITY --(M.OLECULAR AND ATOMIC WEIGHTS MOLECULAR FORMULA MOLECULAR EQUATIONS VALENCE 166 XIV. CARBON AND ITS OXIDES CYANOGEN 181 XV. MF.THANE ETHYLENE ACETYLENE ILLUMINATING GAS ^. FLAME BUNSEN BURNER OXIDIZING AND REDUC- ING FLAME 202 J Contents. HAPTER PAGE XVI. FLUORINE BROMINE IODINE . . . . . . 225 ^+- XVII. SULPHUR AND ITS COMPOUNDS 235 XVIII. SILICON BORON .255 XIX. PHOSPHORUS ARSENIC ANTIMONY BISMUTH . . 265 XX. METALS .278 XXI. SODIUM POTASSIUM LITHIUM 284 XXII. COPPER SILVER GOLD . . . . . .301 XXIII. CALCIUM STRONTIUM BARIUM 319 XXIV. MAGNESIUM ZINC CADMIUM MERCURY . . .331 XXV. ALUMINIUM . . 343 XXVI. TIN LEAD . . . .354 XXVII. CHROMIUM MANGANESE 365 kxVIII. IRON NICKEL COBALT . 373 \XXIX. PLATINUM AND ASSOCIATED METALS . . ..'.. 392 \>J XXX. PERIODIC LAW SPECTRUM ANALYSIS .... 396 4 ,_AXI. SOME COMMON ORGANIC COMPOUNDS 405 APPENDIX ............ 437 PART I DESCRIPTIVE CHEMISTRY DESCRIPTIVE CHEMISTRY. CHAPTER I. INTRODUCTION. CHEMISTRY is a branch of natural science. It deals with the properties of matter, the changes which affect the composition of matter, with numerous laws and theories, and with the manufacture of a vast number of different substances indispensable to the welfare of man- kind. Properties of Matter. Different substances are recog- nized and distinguished by their properties. Color, odor, taste, weight, and solubility are familiar properties ; but to these must be added behavior with heat, light, and electric- ity, and especially the action of different kinds of matter upon each other. Physical and Chemical Changes. Observation shows that the properties of matter can be changed. Sometimes the change is only temporary, as in the freezing of water, or in the melting of iron. Such changes are called physi- cal changes. But often the change is permanent, as in the burning of coal, or the digestion of food. Such changes are called chemical changes. In physical changes the original properties reappear after the cause of the change has been removed. But chemical changes 2 Descriptive ^Chemistry. affect the essential nature of a substance. They are fundamental. Removal of the cause of a chemical change does not restore the original properties of the substance. Thus, coal is readily changed into ashes and invisible gases, but the ashes and gases do not reunite into coal after the heat has been removed. Another essential char- acteristic of chemical changes is the formation of one or more kinds of matter different from the original substance. Thus, water may be decomposed by electricity into two gases hydrogen and oxygen. This is a chemical change, because (i) the water has disappeared, its identity is lost, it has been permanently changed, and (2) other kinds of matter have been formed, which are totally unlike water. Chemistry is largely a study of chemical changes. The different changes which matter undergoes furnish a convenient basis for the classification of properties. Thus, we call physical properties those which accompany physical changes ; while chemical properties require a chemical change for their manifestation. Thus, the color, luster, specific gravity, melting point, and capacity to con- duct electricity are physical properties of copper; but it displays chemical properties when it is heated, or when acted upon by acids, sulphur, and other substances. Examples of simple physical changes are the formation of ice or steam from water, the electrification of a copper trolley wire, the production of colors in the sky, the magnetization of iron in a dynamo or magnet, and \the melting of iron in a foundry. Familiar chemical changes are the rusting of iron, the growth of plants, the burning of oil in a lamp, the decay of fruit, and the souring of milk. Chemical changes are often complex. In many in- stances they are caused by heat, and usually they produce heat. In general, the velocity of chemical change in- creases with rise of temperature. Light induces chemical Introduction. 3 changes, as in growing plants and on photographic plates. Electricity is involved in many chemical changes, a vast industry having recently grown up in this field. Contact is necessary for chemical change, and many substances must be pressed together, intimately mixed, or dissolved before they will interact. Physical and chemical changes are closely related. They usually accompany each other, and are often insep- arable. If the essential change in a substance or sub- stances is chemical, then the substances are said to undergo chemical action. Very often the chemical action involves several substances. The substances are then said to interact or react, and the series of changes is called a reaction. Thus, when zinc is added to nitric acid, the chemical action which occurs is manifested by the forma- tion of a brown gas and the disappearance of the zinc. The zinc and acid interact, and tlie chemical changes can be classified as due to the reaction between zinc and nitric acid. Classes of Chemical Action. There are four general kinds of chemical action, (i) Analysis or decomposition is the separation of matter into its components. Thus, heat decomposes wood, and the juices of our bodies de- compose food. (2) Synthesis or combination is the union of different kinds, or sometimes the same kind, of matter. For example, the gases, hydrogen and oxygen, may be made to unite and form water by passing an electric spark through them. (3) Substitution is the replacement of one kind of matter by another. When zinc is added to hydrochloric acid, the hydrogen leaves the acid, and zinc takes its place. (4) Sometimes parts of different sub- stances exchange places ; this kind of change is called metathesis or double decomposition. If silver nitrate is 4 Descriptive Chemistry. added to hydrochloric acid, the silver and hydrogen ex- change places, forming silver chloride and nitric acid. These four kinds of chemical changes will be fully illus- trated and studied in the succeeding pages. Chemical Energy. We learn in physics that heat, light, and electricity are different forms of energy. They produce special changes. It is also possible to transform the different kinds of energy into each other. Thus, elec- tricity is generated from the heat liberated by burning coal, and electricity in turn may be transformed into light. In chemistry we study another kind of energy, called chemical energy, chemical attraction, or chemism. This is the immediate agent involved in chemical change. Com- bination and decomposition are due to its operation. Chemical energy may be transformed into light, electricity, and heat, and vice versa. Appreciable heat often accom- panies chemical changes, and we shall have many illustra- tions of the intimate relation between heat and chemical energy. Electricity is produced in an electric battery by chemical action. Light is one result of the chemical action called combustion or burning. In fact, every chemi- cal change is accompanied by an energy change of some kind, and in such transformations all the energy can be accounted for, none is lost or gained. Chemical energy is an essential factor in all chemical changes, but we know little or nothing of its nature. We can only study its results and its manner of action. Conservation of Matter. In chemical changes matter is not created or destroyed. It is often transformed, and apparently lost, but the total weight of the substances par- ticipating in any chemical change is always the same. The fact that matter is indestructible was first demon- Introduction. 5 strated by the French chemist, Lavoisier (1743-1794), and countless observers have since shown that it is a funda- mental law of chemistry. The law is called the Law of the Conservation of Matter, and is often stated thus : - No weight is lost or gained in a chemical change. Chemical Elements. Study of the constitution of matter shows that some kinds can be decomposed into substances totally unlike the original matter. Water, for example, is easily decomposed into the gases, hydrogen and oxygen, which are entirely different from water. But it is impossible by any known process to obtain from some kinds of matter substances which have simpler prop- erties than the original substance. Thus, neither oxygen nor hydrogen can be decomposed by any known means. Iron and the familiar metals likewise cannot be divided chemically into two or more substances, nor can they be transformed into each other. They are fundamental sub- stances. We can add other substances to them, but we cannot get simpler substances from them, nor can we transform them into simpler substances. Iron contains nothing but iron. The substances which have such simple properties and at present defy decomposition and trans- formation are called the chemical elements. They are analogous to the letters of the alphabet, and by their vari- ous combinations make up the matter of the universe, some- what as letters form words. There are about eighty elements. Probably there are some undiscovered, but it is generally believed that the present number will not be largely increased. Each element is designated by a symbol, which is an abbreviation of its name. The following is an alphabeti- cal Descriptive Chemistry. TABLE OF THE IMPORTANT ELEMENTS. NAME. SYMBOL. NAME. SYMBOL. Aluminium .... Al Lead .... Pb Antimony .... Sb Lithium Li Arsenic Barium As Ba Magnesium .... Manganese Mg Mn Bismuth Bi Mercury Hg Boron B Nickel Ni Bromine Cadmium Br Cd Nitrogen .... Oxygen N o Calcium Ca Phosphorus p Carbon c Platinum Pt Chlorine Cl Potassium K Chromium .... Cobalt . . Cr Co Silicon Silver Si Ae 1 Copper Cu Sodium "8 Na Fluorine Gold F Au Strontium .... Sulphur Sr s Hydrogen H Tin Sn Iodine I Zinc Zn Iron Fe Of the above elements only eight are abundant in the earth's crust, as may be seen by a TABLE OF THE APPROXIMATE COMPOSITION OF THE EARTH'S CRUST (BY WEIGHT). ELEMENT. Oxygen Silicon Aluminium Iron . Calcium Magnesium Potassium . Sodium Total PER CENT. 47.29 27.21 7 .8l 5.46 3-77 2.68 2.40 2.36 98.98 Introduction. The atmosphere contains about 20 per cent of oxygen and 79 per cent of nitrogen in the free state. The ocean contains about 86 per cent of oxygen, 1 1 per cent of hydro- gen, and 2 per cent of chlorine in combined states. It is clear that the globe, as we know it, is made up of a very few elements. Many of the familiar metals are elements, e.g. lead, zinc, tin, copper, iron, gold, and silver. Other elements besides the metals are solids, such as sulphur, carbon, and phos- phorus ; two are liquid, viz. bromine and mercury ; while several are the common gases, oxygen, nitrogen, and hydro- gen. Many are important simply because they are com- bined with other elements, especially silicon, which is found in most rocks, and calcium, which is a component of limestone. The following is a TABLE OF THE UNCOMMON ELEMENTS. NAME. SYMBOL. NAME. SYMBOL. Ar^on A Prasedymium Pr Beryllium Be Rhodium Rh Caesium Cs Rubidium Rb Cerium Erbium Ce Er Ruthenium .... Samarium Ru Sm Gallium .... Ga Scandium . Sc Germanium Ge Selenium Se Glucinum Gl Tantalum Ta Helium He Tellurium . . . . Te Indium In Thallium Tl Iridium Ir Thorium Th Krypton Kr Titanium Ti Lanthanum La Tungsten W Molybdenum Mo Uranium u Neodymium Nd Vanadium . v Neon . Ne Xenon Xe Niobium Nb Yb Osmium Os Yttrium Yt Palladium Pd Zirconium .... Zr 8 Descriptive Chemistry. Chemical Symbols are usually the first letter of the name of the element. Thus, O is the symbol of oxygen, H of hydrogen, N of nitrogen. Since several elements have the same initial letter, the symbol of some elements contains two letters. Thus, C represents carbon, while the symbol of calcium is Ca, of chlorine Cl, of chromium Cr, and of copper Cu. The symbols of several elements, especially the metals so long known, are derived from their Latin names, as may be seen from a TABLE OF LATIN SYMBOLS. ELEMENT. LATIN NAME. SYMBOL. ELEMENT. LATIN NAME. SYMBOL. Antimony Stibium Sb Mercury Hydrargyrum Hg Copper Cuprum Cu Potassium Kalium K Gold Aurum Au Silver Argentum Ag Iron Ferrum Fe Sodium Natrium Na Lead Plumbum Pb Tin Stannum Sn Symbols always begin with a capital, and are not followed by a period. They should be learned by actual use. Their significance will be explained in later chapters. Chemical Compounds. When elements unite with each other the product of the union is a chemical com- pound. The elements which make up a chemical com- pound are called components. Chemical compounds have three essential characteristics, (i) Their components are held together by chemical attraction. The hydrogen and oxygen, which are the components of water, cannot be separated unless their attraction for each other is over- come by heat, electricity, or some other agent. (2) In any given chemical compound the components are always in Introduction. g the same ratio. Thus, pure common salt, however pre- pared or wherever found, always contains 39.32 per cent of sodium and 60.68 per cent of chlorine. So also water always contains eight parts (by weight) of oxygen and one of hydrogen. Facts similar to these might be given cover- ing all cases examined. Such facts illustrate the general principle that chemical action proceeds according to laws. (3) In chemical compounds the identity of the components is lost. Thus, the red metal, copper, the yellow solid, sulphur, and the invisible gas, oxygen, are the components of the blue solid, copper sulphate. Chemical compounds must not be confused with mixtures. The parts of a mixture may vary in nature and in proportion ; they are also held together loosely, and may often be separated by some mechanical operation, as filtering or sifting. A mixture, too, often has properties similar to its parts. EXERCISES. 1 1. State three properties of (a) glass, (<) wood, (c} water, (W) paper, (e) air. 2. Give three illustrations of (a) physical changes and (6) chemical changes occurring in everyday life. 3. Are the following changes physical or chemical? (a} Burning of wood, () melting of butter, (c) freezing an ice-cream mixture, (d} weathering (i.e. decay) of granite, (e) tarnishing of brass and other metals, (/) formation of snow, (g) developing a photographic plate, (h) seasoning of wood, (/) formation of dew, (/) disappearance of a fog. 4. What ai^ls and what retards chemical change? What often ac- companies it? 5. What physical change accompanies (a} the burning of coal, (6) the action of an electric battery, (c) the burning of a match ? 6. Give an illustration of the transformation of chemical energy into heat, light, or electricity. 7. State the law of the conservation of matter. 1 These exercises are intended for review work. io Descriptive Chemistry. 8. (rt) Name five elements with which you are familiar. () Name the eight most abundant elements in the earth's crust in their order. 9. What common metals are elements? 10. How do elements and compounds essentially differ? Could you prepare (a) a compound from elements, (^) elements from a compound, and (c) elements from elements? 11. Define (a) chemistry, () physical change, (c) chemical change, (d) chemical action, (^) analysis, (/") synthesis, (g) metathesis, (^) sub- stitution, (/) element, (/) compound, () mixture, (/) symbol. 12. Review or learn the metric system (see Appendix, i). PROBLEMS. Perform the problems in the Appendix, i, CHAPTER II. OXYGEN. OXYGEN has played an important part in the develop- ment of chemistry, and is an appropriate element with which to begin a systematic study of this science. Occurrence. Oxygen is the most abundant atid widely distributed of the elements. Mixed with nitrogen and a few other gases, it forms one fifth (by volume) of the atmosphere. Combined with hydrogen, it constitutes eight ninths (by weight) of water; combined with silicon and certain metals, it makes up nearly half of the earth's crust; while compounds of oxygen, carbon, and hydrogen form a large part of animal and vegetable matter. Starch, for example, which is a constituent of all plants, contains about 50 per cent oxygen. Preparation. Oxygen may be prepared from its com- pounds or from air. It was first prepared by decomposing a red compound of oxygen and mercury. When heated in a hard glass tube, this compound decomposes into oxygen and mercury ; the oxygen is collected over water in a pneumatic trough, and the mercury condenses as globules or a film on the upper part of the tube. This experiment is historically interesting, because it was first performed by Priestley, the discoverer of oxygen. The gas is often prepared by decomposing potassium chlorate a compound of oxygen, chlorine, and potassium. Heated to a rather high temperature, the potassium chlo- 12 Descriptive Chemistry. rate passes through a series of changes ; as a final result, the oxygen is set free, and potassium chloride, a white solid, remains behind. Oxygen is most conveniently prepared by heating a mixture of potassium chlorate and manganese dioxide in a glass or metal vessel. The gas is liberated freely from this mixture at a lower temperature than when either compound is heated alone. The manganese dioxide may be recovered unchanged at the close of the experiment. It takes some part in the chemical changes, but just what is not definitely known. It has been suggested that the manganese dioxide combines at first with oxygen, thereby forming another coin- pound of manganese richer in oxygen than the dioxide, but so unstable that when heated it yields oxygen and manganese dioxide. Large quantities of oxygen may be prepared by heating a mixture of potassium chlorate and manganese dioxide in a copper or iron retort. Other commercial processes are used. In Erin's process, which is oper- ated largely in England, purified air is forced by a pump over barium oxide heated to 700 C., 1 thereby forming barium dioxide. The air sup- ply is then cut off, and the pressure in the retorts reduced by reversing the pump. This operation changes the barium dioxide into barium oxide and oxygen. The gas is drawn off into a reservoir. The process is then repeated. A kilogram of barium oxide yields about ten liters of oxygen at a single operation. 2 Oxygen can be prepared from liquid air (see Liquid Air). By evapo- ration at the ordinary temperature and pressure, the nitrogen escapes from the liquid air more rapidly than the oxygen, leaving finally a liquid which is nearly pure oxygen. Unlimited quantities of oxygen may thus be cheaply prepared from the air. This method awaits development. Properties. Oxygen gas has no color, odor, or taste. It is slightly heavier than air. It is somewhat soluble in 1 C. is the abbreviation of " centigrade," which is the name of the thermometer used in science. According to this thermometer water boils at 100 and freezes at o (see Appendix, 2). 2 " Kilogram " and " liter" are denominations of the Metric System of Weights and Measures. This system should be learned or reviewed (see Appendix, i). Oxygen. 13 water, but the presence of even a. small proportion in water is exceedingly important. Fish die in water con- taining no oxygen; and the oxygen absorbed by flowing water helps keep it free from organic matter. (See Decay, below.) The density of oxygen gas is 1.105 (air = i). One hundred liters of water dissolve only about three liters of oxygen under ordinary conditions. The chemical activity of oxygen is its most striking property. It combines with all the other elements except fluorine, bromine, and the inert gases recently discovered in the atmosphere. With most of them the union is direct, and is often accompanied by light and heat, though the temperature at which combination occurs varies between wide limits. At the ordinary temperature it unites with phosphorus, as may be seen by the glow and fumes when the end of a match is rubbed, especially in a dark room. Metals, such as iron, lead, zinc, and copper, tarnish or rust easily, i.e. they combine with the oxygen of the air. The chemical activity of oxygen at high tem- peratures is readily shown by putting burning substances into it. All burn vividly in oxygen. When a glowing stick of wood is put into oxygen, the stick instantly bursts into a flame ; and if left in' the oxygen, the wood continues to burn brightly until the gas is exhausted. If glowing charcoal is put into oxygen, the charcoal burns violently, and throws off showers of sparks. Sulphur burns in air with a small, blue flame, but in oxygen the flame is much larger and brighter. The flame in both cases is accompanied by fumes which smell like a burning sulphur match. Iron wire does not burn in air, but if the end is coated with burning sulphur and then put into oxygen, the wire burns vividly, throwing off a shower of sparks ; when the flame has disappeared, a globule of red-hot iron is often seen on the end of the wire ; and sometimes the inside of the bottle is coated with a reddish powder, which is mainly a compound 14 Descriptive Chemistry. of iron and oxygen. Iron and oxygen combine at a higher tempera- ture than do sulphur and oxygen, so sulphur is used to set fire to the iron. On the other hand, if lighted magnesium is put into oxygen, the burning metal instantly becomes surrounded with a dazzling flame, and burns rapidly to a white powder, thus showing that the temperature at which it combines with oxygen is much lower than that required by iron. Oxidation. When sulphur, iron, magnesium, and car- bon (in wood and charcoal), and other elements burn in oxygen, they combine with it. This chemical change is called oxidation. The fact that oxidation is merely a combining with oxygen may be easily verified. It has been repeatedly shown that oxygen is one con- stituent of all the products formed by burning substances in that gas. Thus, carbon forms an invisible gas called carbon dioxide, which is a compound of carbon and oxygen. Similarly, sulphur, iron, and magne- sium form compounds of these elements and oxygen. These facts may be further verified by a simple experiment. If mercury is heated, it gains in weight, and red particles collect on its surface ; but if it is pro- tected from the air by some coating and then heated, there is no gain in weight and no evidence of the red product. Therefore, when the exposed mercury is heated, something from the air must be added to it. Now, if the red substance is collected and heated in a glass tube, mercury and oxygen are the only products. Hence, the exposed mercury, when heated, must have combined with the oxygen of the air. Oxidation is not always rapid enough to produce light and appreciable heat. Iron and other metals rust, and wood decays slowly, but both processes are mainly oxida- tion. Sometimes oxidation develops considerable heat. Thus, oily rags, piles of hay, and heaps of coal often take fire unexpectedly because of the continued oxidation. Such oxidation is often called spontaneous combustion. Substances which give up oxygen readily are called oxidizing agents. Potassium chlorate is used in fireworks for this purpose, and potassium nitrate acts similarly in gunpowder. In the process of oxidation, oxidizing agents Oxygen. 15 lose oxygen, and are said to undergo reduction a process which will be more fully described in the next chapter. Oxides are formed when oxygen combines with other elements. There are many oxides, and their names express in a general way their composition. Oxides of different elements are distinguished by placing the name of the ele- ment (or a slight modification of it) before the word oxide, e.g. magnesium oxide, lead oxide, zinc oxide. Sometimes di-, or a similar numerical syllable, is prefixed to the word oxide, e.g. carbon dioxide, manganese dioxide, sulphur trioxide, phosphorus pentoxide. The significance of the prefix is explained in Chapter VII. Combustion, in a narrow sense, is rapid oxidation, which is always accompanied by light and heat. Popularly, com- bustion means fire or burning, and substances which burn easily are called combustible. Oxygen is essential to ordi- nary combustion, and is often called a supporter of com- bustion. Exclude air from a fire, and the fire goes out. When coal or wood burns, the carbon (of which they largely consist) unites with the oxygen of the air, forming thereby the invisible gas carbon dioxide,, and the chemical change is manifested by heat and light. /Chemically speak- ing, a substance burning in the air is Uniting rapidly with oxygen. But since the air is about one fifth oxygen and four fifths nitrogen, a gas which does not support com- bustion, it follows that combustion is more vigorous in oxygen than in air. The correct explanation of fire, burning, and combustion was first made by Lavoisier (1743-1794). For many years chemists had be- lieved that all combustible substances contained a principle called phlogiston, and that when a substance burned, phlogiston escaped. Very combustible substances were thought to contain much phlogiston, and incombustible substances no phlogiston. This theory of combus- 1 6 Descriptive Chemistry. tion was proposed by Becher (1635-1682) and advanced by Stahl (1660-1734). Many famous chemists Priestley, Scheele, and Caven- dish supported it. Lavoisier, in 1775, proved by his own and others 1 experiments, that phlogiston did not exist, and that combustion is a process of combination with " a certain substance contained in the air." Soon after he identified this substance as oxygen. The theory of phlogiston, in spite of its falsity, exerted a wholesome influence on the development of chemistry. Combustion, in a broad sense, is not necessarily oxida- tion, but chemical action which develops enough energy to produce light and heat. This broader meaning will be discussed later. Relation of Oxygen to Life. Oxygen is essential to all forms of animal and plant life. If an animal or a plant is deprived of air, it dies. By respiration air is drawn into the lungs and there it gives up part of its oxygen to the blood. This oxygen, which is distributed to all parts of the body by trie blood, oxidizes food and the tissues of the body. As a result of this oxidation new tissue is built up and waste products are formed. One of these waste prod- ucts is carbon dioxide gas, which with other gases is exhaled from the lungs. The blood during its circulation turns dark red, owing to the loss of oxygen ; and when this dark red blood reaches the lungs, it receives a fresh supply of oxygen which turns it bright red, thus preparing it for another journey through the body. Food must be oxidized before it can be taken up by the body, and by this oxida- tion the carbonaceous matter of the body is slowly burned to carbon dioxide. It is this slow oxidation which keeps the body warm. The human body resembles a steam engine. In^each, the oxYggiLjQ-the-air.lielps_burn fuel ^ largely composed of carbon. In the engine, the products es"cape through a chimney and the heat produced is used Oxygen. 17 to form steam which moves parts of the machine ; in the body, the products escape mainly through the lungs and the heat keeps the body at a temperature at which it can best perform its functions. It was formerly believed that breathing pure oxygen would produce too rapid oxidation in the body and burn up the tissue faster than it could be made. But recent study shows that with proper precautions oxygen may be breathed by a healthy person without producing any harmful effect. The blood apparently absorbs a maximum quantity of oxygen, whether supplied from air or from the pure gas. Oxygen is often administered to a person who has been suffocated, or to one who is unable to inhale enough air, as in cases of croup, asthma, or extreme weakness. It is sometimes used to sustain life where air is impure or rare, as in diving bells and submarine boats, and during balloon ascensions to a great height. Decay is in part oxidation. The oxygen of the air together with water vapor acts upon animal and vegetable matter and slowly burns it up. The decomposition is often begun and hastened by bacteria. The products of decay are numerous, carbon dioxide being one. The oxygen dissolved by water assists in the decay of the impurities constantly flowing into rivers. Similarly, it oxidizes in- jurious vapors and matter in the air, literally burning them up, just as it burns wood in a stove. Hence, running water is more likely to be cleaner than standing or stagnant water, and the air in the open country or at the seashore purer than in the crowded city. Uses of Oxygen. Oxygen for commercial use is stored under pressure in strong iron cylinders. The pure gas has limited use, since air, although it contains about 80 per cent of the inert gas nitrogen, may usually be used in place of oxygen, A mixture of oxygen and hydrogen burned in a suitable apparatus produces an intensely hot flame, which is sometimes used to melt refractory metals and to produce the calcium light (see Oxyhydrogen Blowpipe). 1 8 Descriptive Chemistry. Liquid Oxygen. All gases at a low temperature and under great pressure may be condensed to liquids, and even to solids. Under these conditions oxygen becomes first a pale blue liquid and finally a whitish solid. A small quantity was first obtained in 1877, but now it is prepared by the gallon. It is magnetic, and when a strong electro- magnet is held near its surface, the liquid suddenly "leaps up to the poles and remains there permanently attached until it evaporates." Under the normal pressure (760 mm.) 1 liquid oxygen boils at 181.4 C., and at this temperature its specific gravity is 1. 124 (water i). Discovery of Oxygen. Oxygen was discovered on August i, 1774, by Priestley (1733-1804). He prepared it by focusing the sun's rays upon the red mercury oxide by means of " a burning lens of twelve inches' focal distance." It was independently discovered by Scheele (1742-1786), a Swedish chemist, about the same time. Priestley called the gas dephlogisticated air, because he regarded it as " devoid of phlogiston." Scheele called it empyreal air, i.e. fire air or fire-supporting air, because it assisted combustion. Lavoisier, in 1778, gave it the name oxygen (from the Greek oxus, acid, and^w, the root of a verb meaning to produce), because he believed from his experiments that oxygen was necessary for the production of acids a view now known to be incorrect. Weight of a Liter of Oxygen. The volume occupied by a gas depends upon the pressure and temperature to which it is subjected. The volume expands with rise of temperature or with lowering of pressure, but contracts with fall of temperature or with increase of pressure. In general, if we cool a gas or subject it to a pressure, it shrinks, and if we heat a gas or decrease the pressure 1 This expression means the normal or standard pressure of the atmosphere as recorded by the barometer (see Chapter VI). Oxygen. 19 it is under, it expands. Gas volumes, to be correctly compared, must therefore be at the same temperature and pressure. The normal or standard temperature is zero degrees on the centigrade thermometer, or briefly o C. The normal or standard pressure is the pressure of the atmosphere indicated by the barometer when the mercury is 760 millimeters high, or briefly 760 mm. Under these conditions, which are called standard conditions, a liter of dry oxygen weighs 1.43 gm. It is not usually convenient to measure gases at o C. and 760 mm. So if their volumes are to be studied and compared, it is customary to reduce the observed volume to the volume it would occupy under standard conditions. This reduction is accomplished by applying two laws the Law of Charles and the Law of Boyle. Law of Charles. It has been found by experiment that under con- stant pressure all gases expand or contract equally for equal changes of temperature. More explicitly, a gas expands or contracts ^ of its volume at o C. for every degree through which it is heated or cooled. This means that 273 volumes at p become 274 at i, 275 at 2, 280 at 7, 272 at i, 270 at 3, or 273 + 1 volumes at / (i e. at any tem- perature). This law is not absolutely correct, but its variations from the truth are slight. Suppose we have 10 1. of oxygen at o C., and we wish to know the volume it would occupy at 15 C. The problem is easily solved by stating it as a proportion, thus 273:273+ I5::lo:.r. - The value of .ris the volume required. Conversely, in reducing 10 vol- umes at 15 C. to the volume occupied at o C., the proportion is 273 + I5:27 3 ::io:;r. If the given temperature is below o, the number of degrees is subtracted from 273. Law of Boyle. It has also been found by experiment that under constant temperature the volume of a gas is inversely proportional to lo Descriptive Chemistry. the pressure. This is Boyle's law. It means that doubling the pres- sure halves the volume, and vice "versa. Like the above law, this law is only approximately correct. Suppose we have 10 1. of oxygen at 760 mm., and we wish to know the volume it would occupy at 775 mm. According to the law, the proportion expressing the relation is 760:775::^: 10. The value of x\<& the required volume. Conversely, if we have 10 1. at 775 mm., and wish to know its standard volume, the proportion is It is convenient to notice that the proportion is stated so that the extremes (or means) are the original pressure and volume. In other words, one pressure multiplied by its volume equals the other pressure multiplied by its volume, or P\P\\V\ V. Hence, the proportion is applicable to values not necessarily includ- ing 760. EXERCISES. i . What is the symbol of oxygen ? 2. How is oxygen prepared (a) in the laboratory, and () commer- cially ? 3. Name several compounds from which oxygen can be prepared. 4. Summarize the properties of oxygen. What is its most charac- teristic property ? 5. If air contains something besides oxygen, what must be the gen- eral properties of this other ingredient ? 6. Define and illustrate (a) oxidation, (ft) oxide, (c) combustion, (//) oxidizing agent. 7. What elements were mentioned in studying oxygen ? What compounds ? 8. What general chemical change is involved in burning ? What class of chemical changes is illustrated by (a} preparation of oxygen from mercuric oxide, (b} burning of sulphur in oxygen ? 9. Give a brief account of Priestley, Scheele, and Lavoisier (see Appendix, 4). Oxygen. 21 10. What chemical part does oxygen 'take in (a) respiration, (#) de- cay, (c) combustion, () oxidation ? 11. State and illustrate (a) Charles's law and () Boyle's law. 12. Give a brief account of Boyle and of Charles. PROBLEMS. 1. Potassium chlorate contains about 39 per cent of oxygen. How many grams of oxygen can be prepared from (rt) 100 gm., (^) 250 gm., and (c) 725 gm. of potassium chlorate ? 2. What approximate weight of oxygen can be prepared from 100 gm. of potassium chlorate containing 12 per cent of impurity ? 3. What is the weight of (a) 10 1. of oxygen, (b) 75-!., (c) 500 cc., (d) 750 cc., 0) 4!.? 4. A room 25 m. long, 17 wide, and 15 high is filled with oxygen. What weight of gas does it contain ? (A liter of oxygen weighs 1-43 gm.) 5. Reduce the following volumes to the volume occupied at o C. : (a) 173 cc. at 12 C., (b) 466 cc. at 14 C., (c) 706 cc. at 15 C., (d) 25 cc. at 27 C. 6. A volume of gas at o C. measures 1500 cc. What is its volume at (a} 15 C., (d) 50 C., (0 100 C., (d} 300 C. ? 7. If 500 cc. of gas at 27 C. are cooled to 5C., what is the new volume ? 8. Reduce the following volumes to the volume occupied at 760 mm. : (a) 200 cc. at 740 mm., (b) 25 cc. at 780 mm., (c) 467 cc. at 756 mm. Ans. (a) 1947? (^) 25.65, (c) 464-54- 9. A gas measures 1000 cc. at 770 mm. What is its volume at 530 mm.? 10. Reduce the following to standard conditions: (a) 147 cc. at 570 mm. and 136.5 C., (b} 320 cc. at 950 mm. and 9iC, (c) 480 cc. at 380 mm. and 68.25C, (d) 25 cc. at 780 mm. and 27 C., (*) 14 cc. at 763 mm. and iiC. Ozone is a gas related to oxygen, though its properties differ. It is formed when electric sparks pass through the air, and is therefore pro- duced when electrical machines are in operation and during thunder storms. Slow oxidation, especially of moist phosphorus, produces ozone. Indeed, its formation accompanies several chemical changes. 22 Descriptive Chemistry. such as the burning of hydrogen and of certain resins, and the decom- position of water by electricity. Ozone has a peculiar odor, suggesting burning sulphur. The name ozone signifies smell. It is active chemically, tarnishing metals, bleach- ing colored vegetable substances, deodorizing foul animal matter, and corroding such substances as cork and rubber. It is sometimes used as a disinfectant, though other oxidizing agents are more convenient. When heated to 250 C., or higher, it is wholly changed into oxygen. Ozone, therefore, contains nothing but oxygen. When oxygen is changed into ozone, it is found that three volumes of oxygen yield two volumes of ozone ; and, conversely, the two volumes of ozone, when heated, become three volumes of oxygen. Hence, volume for volume, ozone is 1.5 times heavier than oxygen. For this reason ozone is sometimes called "con- centrated oxygen," or "an oxide of oxygen." Its theoretical relation to oxygen will be subsequently discussed. The atmosphere usually contains a small proportion of ozone, prob- ably not more than one volume in 700,000 volumes of air. It is more abundant in the open country and at the seashore than in cities. CHAPTER III. HYDROGEN. Occurrence. Free hydrogen is present in the gases petroleum wells, and natural gas openings. Artificial illuminating gas contains consid- ^ erable hydrogen. It is^T product of fermentation and *"* decay, and according to recent observations a very small quantity is present in the atmosphere of the earth. Enor- * mous quantities of free hydrogen exist in tne atmqsrjhere 7 of the sun, and during an eclipse of the sun gigantic streams of burning hydrogen may be seen shooting out from the sun's disk thousands of miles into space. Other heavenly bodies which are self-luminous, like the star Sirius and the nebulae, contain free hydrogen. The spectroscope has revealed its presence in these distant bodies. Meteor-*? ites^ which come from regions far beyond our earth, often contain free hydrogen. Cojnbinedji^drogen is abundant and widely distributed. * It forms one ninthly weight^ of water. Most animal and f ^ vegetable matter contains hydrogen. It is also an essential 1 1 component nf_a1J_gHHs Combined with carbon, it forms many gases and liquids called hydrocarbons, which are con- ' * stituents of illuminating gas, kerosene, and naphtha. Com- bined with carbon and Oxygen, It forms many vegetable > C> compounds, such as sugar, starch, parser, wood, and numer- ous artificia'1 products. With nitrogen it forms the familiar ^ compound, ammonia ; and with sulphur, the bad-smelling gas, ^ hydrogen sufphTdeTwhich occurs in many sulphur springs. ' 23 24 Descriptive Chemistry. Preparation. Hydrogen, like oxygen, is prepared from its compounds. In the laboratory this is easily accom- plished by allowing a metal and an acid to interact. The metals usually employed are zinc, iron, or magnesium, and the acids are dilute sulphuric acid or hydrochloric acid. The hydrogen comes from the acid and bubbles through the liquid, when the acid and metal are put into a test tube or flask. On a large scale hydrogen is prepared in a genera- tor, which consists of a glass vessel provided with a delivery tube arranged to collect the gas over water in a pneumatic trough. No flame should be near during the performance of this experiment, because mixtures of air and hydrogen explode violently when ignited. The interaction of zinc and sulphuric acid produces, besides hydrogen, a compound called zinc sulphate. This remains in the generator in solution, and if the solution is allowed to evaperate, the zinc sulphate separates as transparent crystals, which soon turn white in the air. Hydrogen may be obtained from water by allowing tiie_Jiilal-ao^lijLir^^ to interact. If a small piece of sodium is dropped upon cold water, the sodium melts into a shining globule, which spins about rapidly on the water with a hissing sound, and finally disappears with a slight explosion. But when the sodium is wrapped in a piece of tea lead pierced with a few holes and then dropped beneath the shelf of a pneumatic trough filled with water, the action proceeds smoothly. Hydrogen gas rises and displaces the water from a test tube or bottle supported over the hole in the shelf. The nature of the chemical change which attends the liberation of hydrogen from water will be explained later (Chap- ter V). Hydrogen, together with oxygen, is liberated from water by passing a current of electricity througlTwafer containing a little sulphuric acid (see Chapter V). Hydrogen may also be prepared by passing steam the gaseous form of water over heated metals. Hydrogen. 25 This experiment was first performed by Lavoisier, in 1783, while he was studying the composition of water. He -passed steam through a red-hot gun barrel containing bits of iron. The oxygen of the steam combined with the iron, and the hydrogen escaped from the tube. Since Lavoisier was studying the composition of water, and not the properties of hydrogen, he naturally thought of this gas as essential for forming water. So he says in his notes, " No name appears to us more suitable than that of hydrogen, that is to say, 'generative principle of water.'" Apart from historical interest, this experiment has commercial value. If steam is passed over red-hot coal (instead of iron), producer gas is formed. This is a mixture consisting largely of hydrogen, which is used as a source of heat in making steel and glass. If oil vapor is added to this mixture, water gas is formed. This is an illuminating gas like ordinary illuminating gas, and is used in many cities (see Water Gas). Physical Properties. Hydrogen has no taste or color. The pure gas has no odor, though hydrogen as ordinarily prepared has a disagreeable odor, due mainly to impurities in the metals used. Most of these impurities may be re- moved by passing the gas through a solution of potassium permanganate. Hydrogen is-the lightest known substance. One liter of dry hydrogen at o C. and 760 mm. weighs only 0.0896 gm. Volume for volume, air is about 14.4 times, oxygen 16 times, and water 11,000 times heavier than hydrogen. The extreme lightness of hydrogen may be easily shown, (i) If a wide-mouth bottle of the gas is left uncovered two or three minutes and a lighted match then dropped in, the match will continue to burn. If hydrogen had been present, the flame would have caused it to combine with the oxy- gen of the air with a loud FlG le l pouring hydroge n. explosion. (2) If a bottle of hydrogen is held beneath a bottle of air as shown in Figure i, the gases 26 Descriptive Chemistry. soon exchange places, the hydrogen, owing to its lightness, rising into the upper bottle. Its presence there may be readily shown by dropping a lighted match into this bottle ; if the experiment has been well done, the hydrogen will burn, but in most cases the loud explosion shows that only a part of the hydrogen has been poured upward. A lighted match dropped into the other bottle reveals only air. (3) If a small collodion, or rubber, balloon is filled with hydrogen and then released, it will rise rapidly into the air. Hydrogen, because of its lightness, is sometimes used to fill large balloons, but ordinary illuminating gas is usually employed. Hydrogen is the standard for reckoning the density of gases. Thus, since a liter of oxygen weighs 1.43 gm., its density is found by the proportion:- Q ^ . , ^ . . , . ^ . ^ l6 Hydrogen is not very soluble in water, but it is absorbed by several metals, especially the rare metal palladium. This property of absorbing gases is called occlusion. Only about 1.84 1. of hydrogen at 760 mm. pressure dissolve in 100 1. of water at 20 C. Palladium absorbs from 370 to 960 times its own volume of hydrogen, according to the conditions of the experi- ment. Platinum and iron act similarly, though to a less degree. Illu- minating gas, which contains considerable hydrogen, is also absorbed by metals. And since heat is developed by occlusion, the illuminating gas may be lighted by the heated metal upon which it flows. A self- lighting gas burner acts on this principle. The act of occlusion is partly chemical and partly physical. Hydrogen illustrates diffusion; i.e. it readily passes through porous substances and completely mixes with other gases without stirring or agitating. It penetrates unglazed earthenware, paper, and heated metals, espe- cially platinum. Hydrogen has the highest rate of diffusion, because its density is the lowest. The rate of diffusion of a gas is inversely proportional to the square root of the density. Thus, the rate of diffu- sion of hydrogen is four times that of oxygen, since the density of oxy- gen is sixteen times that of hydrogen. We are largely indebted for our knowledge of diffusion to the English chemist, Thomas Graham (1805-1869). Hydrogen. 27 Hydrogen is not poisonous if pure. It does not sup- port life, but a little may be breathed without danger. When the lungs are filled with it the voice becomes very shrill and thin. Chemical Conduct. Hydrogen burns in the air and in oxygen with an almost invisible but very hot flame. Water is the product of its combustion. These facts may be verified by the apparatus shown in Figure 2. The hydro- gen, which is generated from zinc and hydrochloric acid in the flask, passes through the U-tube filled with calcium chloride (to remove the mois- ture), and is lighted at the tip after it has driven all the air from the apparatus. 1 A platinum or copper wire held in the flame instantly becomes red-hot. If a small, dry, cold bottle is held over the flame, moisture is deposited inside the bottle. The film of water often noticed on the bottom of a vessel placed over a lighted gas range or a Bunsen burner is formed by the burning hydrogen and hydrogen compounds of the illuminating gas. Similarly, water often drops from the top of the oven of a lighted gas range. Or- ganic substances containing hydrogen, such as wood and paper, when burned, yield water as one of their products. The fact that the only product of burning hydrogen is water was first shown in 1783 by Cavendish (1730-1810). Lavoisier in the same year verified this fact and utilized it to explain the composition of water. The temperature of the hydrogen flame is very high. More heat is produced by burning hydrogen in oxygen FIG. 2. Apparatus for burning hydrogen. . 1 This experiment is dangerous. The precautions to be observed can be found on pages 48-49 in the author's " Experimental Chemistry." 28 Descriptive Chemistry. than by burning the same weight of any other substance (see Chapter X). Hydrogen burns in chlorine gas. The flame is bluish white, not very hot, and the product is hydrochloric acid gas a compound of hydrogen and chlorine. This burning of hydrogen in chlorine illus- trates the broader use of the word combustion, since no oxygen is involved. Hydrogen does not support combustion, as the term is usually used. This fact is illustrated by putting a lighted taper into an inverted bottle of hydrogen. The taper ignites the hydrogen, which burns at the mouth of the bottle. The taper does not burn inside the bottle, but when it is slowly withdrawn through the burning hydrogen it is relighted. Hence, hydrogen burns, but does not support combustion. A mixture of hydrogen and air explodes violently when ignited. Therefore, the air should be fully expelled from the apparatus in which hydrogen is being generated before the gas is collected, and no flames, large or small, should be near. Neglect of these precautions has caused serious accidents. Hydrogen not only combines energetically with frea oxygen, but it withdraws oxygen from compounds. As stated before, this chemical removal of oxygen is called reduction. Hydrogen is a vigorous reducing agent. The Oxyhydrogen Blowpipe utilizes the intense heat pro- duced by burning a mixture of hydrogen and oxygen. The apparatus (Fig. 3) con- sists of two pointed metal tubes. The inner and smaller one is for the Blowpipe tip. oxyg e n , an d the outer and larger one for the hydrogen. Their pointed ends are Hydrogen. 29 close together, and the two gases mix as they are forced out of these small openings by the pressure maintained in the storage tanks. Sometimes the tubes are separated, but the gases flow from a similar opening. The hydrogen is first turned on and lighted at the pointed opening ; then the oxygen is turned on and the flow gradually regulated until the flame is the desired size, usually thin, straight, and as long as the apparatus requires. There is no danger in using the blowpipe, provided it does not leak and the pressure is properly regulated by the stopcocks. In the hot flame, some metals, like silver, turn to vapor ; some, like iron, burn brilliantly ; while others, like platinum, melt. When the flame strikes against a piece of lime of other sub- stance difficult to melt, the lime becomes intensely bright. Thus used, it is called the lime, calcium, or Drummond light and is often employed in operating the stereopticon. The blast lamp is a modification of the oxyhydrogen blowpipe. The apparatus (Fig. 4) consists of two tubes, an inner one for air and an outer one for illuminating gas. The air, which is forced through the apparatus by a bellows, provides oxygen, and the illumi- nating gas contains hydrogen and other combustible gases. The mixture burns at the opening of the tubes with a colorless or bluish flame, which is hotter than the Bunsen flame the usual source of heat for chemical experiments. The shape of the flame is easily regulated by stopcocks. Liquid Hydrogen is a colorless, trans- parent liquid produced bv subjecting the FIG. 4. Blast lamp, gas to great pressure and low temperature. It was first produced in 1898 by Dewar. The temperature used was 205 C., and the pressure was 180 atmospheres (i.e. 180 times 760 mm.). At the ordinary pressure it boils at 238 C. Under reduced pressure and at 256 C. it becomes "a white mass of solidified foam." jo Descriptive Chemistry. Discovery of Hydrogen. Paracelsus in the sixteenth century ob- tained hydrogen by the interaction of acids and metals. It was iden- tified as an element in 1766 by Cavendish, who called it inflammable air. The name hydrogen, given to it by Lavoisier, in 1783, is derived from the Greek words hudor, water, and gen, the root of a verb mean- ing to produce. EXERCISES. 1. What is the symbol of hydrogen ? 2. What familiar compounds contain hydrogen? 3. How is hydrogen prepared in the laboratory? Describe other methods of preparation. 4. Summarize the properties of hydrogen. What is its most char- acteristic property ? 5. Why is there danger of an explosion in generating hydrogen? How may the danger be avoided ? 6. What is the weight of a liter of dry hydrogen? How many times heavier than a liter of hydrogen is one of air ? 7. Define and illustrate (a) occlusion and (b} diffusion of gases. 8. What chemical change occurs when hydrogen burns in air ? 9. Is water an oxide ? Why ? 10. How does the heat of the hydrogen flame compare with its luminosity ? n. Define (#) reduction and () reducing agent. Name a reduc- ing agent. 12. Describe () the compound blowpipe and (&) the blast lamp, and state the use of each. 13. Summarize briefly the discovery of hydrogen. Give a short account of Cavendish. Why and by whom was hydrogen so named ? 14. What class of chemical changes is illustrated by () the prepara- tion of hydrogen from zinc and sulphuric acid, (<) the burning of hydrogen in air ? PROBLEMS. 1. How many times heavier than a liter of hydrogen is a liter of oxygen, both being dry and under standard conditions ? 2. What is the weight of (a) 500 cc. of dry hydrogen gas at o C. and 760 mm. ? (b) Of 1800 cc. ? (V) Of 9 1. ? 3. The standard pressure at which a gas is measured is 760 mm. Express the same in inches. CHAPTER IV. GENERAL PROPERTIES OF WATER. WATER is worthy of extensive study because of its importance in the animal, vegetable, and mineral king- doms, its peculiar properties, and its numberless uses. Occurrence in Nature. Water, in the form of vapor, is always present in the atmosphere. Evaporation is con- stantly taking place from the surface of the ocean, from the moist earth, from the bodies of animals, and from plants. This vapor is continually condensing, and appears as clouds, mist, fog, rain, snow, hail, dew, and frost. The proportion of water vapor in the atmosphere varies between wide limits, the amount present being largely influenced by the temperature. It has been found, however, that 1000 volumes of ordinary air contain about 14 volumes of water vapor. The total amount of vapor in the atmos- phere is beyond comprehension. In the liquid state water occurs in vast quantities. About three fourths of the surface of the globe is covered with water. Soil and porous rocks hold considerable quantities, and plants and animals contain a large pro- portion. Many substances which are apparently dry really contain a large proportion of water. Thus, in a ton of clover hay there are upwards of 200 Ib. of water, and a ton of salt hay, which is usually very dry, contains about 100 Ib. Many common foods are largely water, as may be seen by the following 3' Descriptive Chemistry. TABLE OF THE PROPORTION OF WATER IN FOOD. FOOD. PER CENT OF WATER. FOOD. PER CENT OF WATER. Cod .... 8^.6 Q4..3 Beef . . 6l.Q Apples 84.6 Lobster "'y 7Q 2 Strawberries . . QO A. Ecrorg . /y * T\-1 Watermelon .... y w "4- 02.4. Asparagus 04.. Milk 87. Potatoes 78.7 Cheese ....... 28 to 72 Cucumbers 954 White bread .... 35-3 The human body is nearly 70 per cent water, and during a year the average man drinks about half a ton. Water in the form of ice permanently covers the coldest parts of the surface of the earth, e.g. the polar regions and the summits of high mountains. A rough estimate of the total weight of ice on the earth's surface is 6,373,000,0x30 millions of metric tons. 1 Functions of Water in Nature. Since water is the only liquid occurring in large quantities on the earth's sur- face, it is the great agent of erosion. It cuts away the earth's crust, and transports the material from higher to lower levels, or washes it into the ocean. Together with carbon dioxide gas it decomposes the rocks, changing them into clay, sand, and substances which make the soil pro- ductive. Its cycle of changes from liquid to vapor and vapor to liquid exerts a marked influence on the distribu- tion of heat and moisture upon the earth's surface, i.e. on climate. It dissolves many solids and gases and is constantly re- moving from the rocks and soil their soluble constituents, 1 A metric ton contains 2204.6 pounds. Properties of Water. 33 some of which serve for the nutrition of plants, though the larger part passes on to the ocean. The latter thus be- comes a vast reservoir of water containing salt and other mineral matter obtained from the earth's crust. In the vital processes of animals and plants it helps change the food into a condition fit for distribution and assimilation. Industrial Applications. Besides the universal use of water for drinking, it is applied to an endless variety of use- ful and convenient purposes. It has always been man's beast of burden. It is the vehicle for transferring mechan- ical energy to water wheels an application now being made on a vast scale for generating electricity. It utilizes by its peculiar properties the energy in fuel by means of the steam engine. It is the highway for transportation on the largest scale by ocean, river, lake, and canal. It is the vehicle for the distribution of heat by hot water and steam. It is the indispensable solvent in metallurgy, in the manu- facture of chemicals, and in such industries as soap making, bleaching, brewing, dyeing, and tanning; it is necessary wherever mortar and cement are used. Man's work would be stopped in a thousand other ways were he deprived of water. Physical Properties of Pure Water. Owing to its remarkable solvent power, water is never found pure in nature, and is purified even in the laboratory only by taking especial precautions. At the ordinary temperature water is a tasteless and odorless liquid. It is usually colorless, but thick layers are bluish. Water is a poor conductor of heat. This last property may be shown by boiling water near the surface in a large test tube containing a piece of ice weighted down upon the bottom. The ice remains unmelted for some time, although the water is boiling a few inches above it. 34 Descriptive Chemistry. Most liquids expand with heat and contract with cold. Water is an exception. If water at 100 C. is gradually cooled, it contracts in volume. But when 4 C. is reached, if the cooling continues, the volume increases as long as the liquid state is maintained. Hence at 4 C. a given volume contains the greatest weight of water. That is, water has its maximum density at 4 C. The density of water at 4 C. is i ; and water at this temperature is the standard for determining the densities of solids and liquids. Thus, when we say the density of gold is 19, we mean that gold is 19 times heavier than an equal volume of water at 4C. The expansion of water when cooled from 4 C. to o C. is slight, but the change is exceedingly important in nature. When the water on the surface of a lake or river cools, it contracts, and since it is heavier (volume for volume) than the warmer water beneath, it sinks. The warmer water rises, is cooled, and likewise sinks, thus causing a circula- tion which continues until all the water from surface to bottom has the temperature of 4C. Now if the cooling continues, the surface water expands and remains on top, because it is lighter than the water beneath. Hence when the temperature of the air falls to oC, this top layer of water freezes and protects the remaining water from the cold, thus stopping the circulation. Should the circulation continue, as the temperature fell from 4 C. to o C., the whole body of water would finally freeze from top to bottom. This condition would not only destroy the fish and marine plants, but seriously affect climate, since the heat of summer could not melt such a vast mass of ice. When water freezes, it expands about one tenth of its volume. That is, 100 cc. of water produce about no cc. of ice. In other words, 100 cc. of water and 1 10 cc. of ice weigh 100 gm. Hence ice floats. The specific gravity of ice is about 0.92. The pressure exerted by water when it freezes is powerful. Vessels or pipes completely filled with water often burst when the water freezes. It is an erroneous but popular idea that " thawing out " a pipe bursts it. As a matter of fact, ice contracts when it melts. The pipe cracks when the water freezes, and as the ice melts a channel is left for the water to Properties of Water. 35 flow out of the pipe. Because of this property, ice is an effective agent in splitting rocks. Water creeps into cracks, especially into the narrow ones by capillary attraction, and when it freezes, the rock is slowly split apart. Water in freezing also destroys the tissue of living plants, which are often said to have been "touched by frost." Frozen flesh for a similar reason becomes pulpy and is more liable to putrefy when thawed a fact sometimes overlooked by those who eat flesh food which has been kept in cold storage. FIG. 5. Snow crystals. From photographs by Wilson A. Bentley. Ice melts at o C. (32 F.), which is also the freezing point of water. Ice often crystallizes in freezing, but the 36 Descriptive Chemistry. individual crystals are seldom visible except during the first stages of the process. Snow crystals are common (Fig. 5). They are always six-sided, and are formed in the atmos- phere by the freezing of water vapor. Water evaporates at all temperatures, passing off as an invisible vapor into the atmosphere or into the air confined over it. If water is heated, the vapor passes off rapidly until the thermometer reads iooC. (or 212 F.). At this point water boils, i.e. it changes rapidly into vapor without rise of temperature. This vapor, if allowed to escape into the atmosphere, cools and condenses quickly into a cloud of minute drops of water. This cloud is popularly called steam. Scientifically, steam is invisible. What we call steam is a mass of very small particles of water. This may be illustrated by boiling water in a large flask. The inside of the flask is perfectly transparent, although there is a cloud of " steam " issuing from its mouth. Water boils when its vapor escapes with sufficient pressure to over- come the pressure of the atmosphere upon its surface. Hence the boil- ing point depends upon the pressure either of the atmosphere or of the vapor within the vessel. The boiling point is iooC. (or 2I2F.) when the atmospheric pressure is normal, i.e. 760 mm. The boiling point is lower as the pressure is decreased and higher as the pressure is increased. Warm water will boil under the receiver of an air pump or on the top of a high mountain. In the city of Mexico (7500 feet above sea level) water boils at about 92 C., and in Quito in South America (9350 feet above sea level) water, which boils at about 90 C., is not hot enough to cook potatoes. The pressure which water vapor exerts as it escapes from a liquid is called its vapor tension. Since the rate of evaporation depends upon the temperature of the liquid, vapor tension varies with the temperature. Vapor tension is usually expressed in millimeters of mercury. Thus, at iooC. the vapor tension of water is 760 mm., because at the boiling point the vapor pressure is just enough to overcome the opposing atmospheric pressure. At 20 C. the vapor tension of water is 17.39 mm - Properties of Water. 37 A liter of steajn, if it could exist at oC. and 760 mm. pressure, would weigh 0.806 gm., or nine times more than a liter of hydrogen. Natural Waters. Water is never found pure in nature. Even rain water, which is usually regarded as the purest natural water, contains gases and dust washed from the air. When rain strikes the ground it begins at once to take up impurities from the rocks, soil, and vegetation. Some of the water flows along the surface, becoming more and more impure, and finally reaches the ocean. From 25 to 40 per cent of the annual rainfall in temperate regions soaks into the ground and trickles through the soil at an estimated rate of 0.2 to 20 feet a day. This underground water finally finds its way again to the surface as a spring or well, through a lake or river, or from a hillside. On its journey underground the water loses most, often all, of its organic matter, remnants of vegetable and animal matter, but dissolves mineral matter and gases. If the amount of dissolved matter in spring water is large or the kind of matter is so unusual as to give the water a marked taste or medicinal properties, the water is called mineral water. Water containing calcium and magnesium compounds is hard, but in soft water, such as rain water, these com- pounds are absent. There are several hundred mineral springs in the United States. Those having a high temperature are called thermal, as at Hot Springs, Arkansas, and at Bath, England. Many contain a large proportion of common salt, as at Saratoga, New York. Others contain alkaline matter and carbon dioxide gas, eg. Vichy and Apollinaris water. Sulphur springs contain solid or gaseous compounds of sulphur or both and have valuable medicinal properties. Some, like Hunyadi, are bitter; but others, especially those in New York State, which contain gaseous sulphur compounds, have a sweet taste but an unpleasant odor. Cha- lybeate waters contain soluble iron compounds. Many waters contain lime and magnesium compounds, and a few contain alum. Most natural Descriptive Chemistry. mineral waters contain traces of a large number of different substances. Many commercial mineral waters have doubtful medicinal value. River water obviously contains the impurities brought by springs and the surface water ; it is also often made very impure by decaying animal and vegetable matter, which has been purposely or accidentally introduced, espe- cially if the river passes through a thickly settled region. A sluggish river is more apt to be impure than a swift one, because the latter tends to purify itself by exposing its impurities to the oxidizing power of the air. Ocean water contains a large proportion of common salt. The other substances in order of abundance are magnesium chloride, magnesium sulphate, calcium sulphate, and potas- sium chloride ; many other substances are present in small quantities. The peculiar taste of ocean water is due to the presence of these substances, and since by evaporation the water only is removed, the ocean always has a " salty " taste. The composition of some natural waters is sum- marized in the following TABLE OF COMPOSITION OF NATURAL WATERS. SOLIDS PARTS PER 100,000. GASES CUBIC CENTIME- TERS PER LITER. KINDS OF WATER. Organic Matter. Calcium Com- Magne- sium Com- Com- mon Total Residue. Nitro- gen. Oxygen. Carbon Dioxide. pounds. pounds. Salt. Rain . I 5 3-4 I3-I 6. 4 i-3 River(Thames) 34 20 1.8 2.6 29 15 7-4 30-3 Spring . . Trace 2 20 I 5 .8 8.6 i Mineral (Bath) Trace 137 23 34 236 4 2 29 Ocean . . Trace 140 530 2650 3500 12. 1 6 17 Properties of Water. 39 Drinking Water. " Water used for drinking should be free from visible suspended particles, without disagreeable taste or smell, and not capable of acquiring such by standing for a day or two in a clean, well- closed vessel. It should also contain enough of the gases derived from the atmosphere to give it a slight taste distinguishable from the flatness of boiled or distilled water. It should not contain solid matter in solu- tion to the extent of more than 300 parts in a million, or about 3 gm. to 10 1. The mineral portion of this solid matter should not con- tain any poisonous substance. As little as possible of the solid contents should consist of organic matter usually not exceeding 15 to 20 parts per million, or about 2 gm. to 100 1. And it is particularly de- sirable that decomposing animal matter or substances which give evi- dence of its previous presence should be found, if at all, as a mere trace. Above all, drinking water should be free from disease-producing bacte- ria or other injurious microorganisms." The problem of obtaining drinking water in large quantities is usually local. In some cities the water is purified by filtering it through a layer of sand and gravel, an acre or more in area and several feet deep. Such a filter removes bacteria almost completely, though it must be frequently cleaned. Sometimes the water is stored in a large settling basin or reservoir and purified by adding alum, or a similar substance, which causes the suspended matter to settle. Dissolved substances cannot be removed without considerable difficulty, so as a rule water is taken from a source which is reasonably pure. The purity of drinking water is usually determined by a water analy- sis. This is not a decomposition of water, but a chemical examination of a sample for the presence and amount of certain substances which indicate or cause impurity. A chemical examination is of limited value, however, unless it is supplemented by a microscopic study of a fresh sample and a rigid sanitary inspection of the premises. Water which is clear, sparkling, cool, attractive to the eye, and pleasant to the taste may be seriously polluted by disease germs, or may be liable to sudden contamination from some unsuspected source. On the other hand, a rather unpleasant-looking water may be harmless. Hence the necessity of careful and extended examination of water to be used as a beverage. Water may be purified by distillation. This operation is not con- venient with large quantities. It is performed in the laboratory in a condenser, which is shown in Figure 6 arranged for use. The condenser consists of an outer tube, A A, provided with an inlet 4 o Descriptive Chemistry. and an outlet for a current of cold water, which surrounds the inner tube, BB. The vapor from the water boiling in the flask, C, condenses FlG. 6. Condenser arranged for the distillation of water. in the inner tube, owing to the decrease in temperature, and drops off the lower end of this tube, as the distillate, into the receiver, D, while the impurities remain behind in the flask. Distilled water is prepared on a large scale in metal vessels, and the vapor is con- densed in a block tin pipe coiled around the inside of a vessel through which a current of cold water is flow- ing. This coiled pipe is called a worm (Fig. 7). Distilled water is used in the chemical laboratory ; large quantities are made into ice. Distillation is an old process. A quaint still is shown in Figure 8. Dis- tillation is the process used to separate liquids from solids and from each other, FIG. 7. Worm- and finds extensive appli- FIG. 8. A quaint still, shaped tube. cation in the manufacture of liquors and kerosene oil. Properties of Water. 41 Solution. Many solids, liquids, and gases disappear when put into water. This operation is called dissolving, or putting into solution. The resulting liquid is called a solution of the substance used. The liquid in which the substance dissolves is called the solvent, and the dissolved substance is called the solute. If the solute is not vola- tile, or not very volatile, it may be recovered by evaporat- ing, or distilling off, the water. The degree of solubility is usually expressed by the terms sligJitly soluble, soluble, and very soluble. It is more accurate, and usually desir- able, to state the proportions of solvent and solute, and also the temperature. Thus, instead of saying that common salt is very soluble in cold water, it is better to state that 36 gm. of salt dissolve in 100 cc. of water at 20 C. Substances which do not dissolve in water are called insoluble, though this term is also applied to those substances a minute quantity of which dissolves in water. Thus glass, sand, and many rocks are usually classed as insoluble substances, but they dissolve appreciably in water. A solution which contains a small proportion of solute is called a dilute solution ; one containing a large propor- tion is called a concentrated solution. Thus, dilute sul- phuric acid usually contains one volume of acid to three or more volumes of water, while concentrated sulphuric acid is nearly 98 per cent acid. Sometimes the terms weak and strong replace dilute and concentrated, but they are ambiguous, and their use should be avoided. Solutions of Gases. Water dissolves or absorbs many gases. The degree of solubility depends upon the gas, the temperature of the water, and the pressure at which solution occurs. Some gases, such as ammonia and hydro- chloric acid gas, are very soluble in water. Advantage of Descriptive Chemistry. this fact is taken in manufacturing ammonium hydroxide and hydrochloric acid. Each commercial substance is merely a water solution of the respective gases, ammonia and hydrochloric acid gas ; the gas is readily liberated by heating the liquid. The common gases, oxygen and hydrogen, are only slightly soluble in water. Air dissolves in water, as may easily be shown by heating faucet water, bubbles of air forming and escaping quickly as heat is ap- plied. Carbon dioxide gas is very soluble in water. Water containing this gas is called "soda water," or carbonated water. More gas is forced into the water than will dissolve at the ordinary temperature and pres- sure, as may be seen by the rapid escape of gas when the water is drawn from a soda fountain. This rapid escape of a gas is called efferves- cence. " Soda water " must, therefore, be stored in a strong vessel and kept in a cool place. The gas was formerly obtained from sodium bi- carbonate a compound related to "soda' 1 ; hence the name "soda water. 11 It is now prepared from marble and an acid, or from liquid carbon dioxide. The volume of gas which will dissolve in water decreases with rise of temperature. Thus, 100 cc. of water at oC. will dissolve 179.6 cc. of carbon dioxide, but only 90.1 cc. at 20 C. The volume of a mod- erately soluble gas which is dissolved by water is directly proportional to the pressure if the temperature is constant. This is Henry's law. It is illustrated by the following TABLE OF SOLUBILITY OF CARBON DIOXIDE GAS. VOL. OF WATER AT o C. VOL. OF CARBON DIOXIDE MEASURED UNDER NORMAL CONDITIONS. PRESSURE' IN ATMOSPHERES. I 1. 900 cc. 1800 cc. 3600 cc. 5 i 2 7200 cc. 4 The tremendous pressure to which subterranean gases are subjected accounts for their presence, especially carbon dioxide, in such large pro- portions in the waters of mineral springs. Properties of Water. 43 Solutions of Liquids. The solubility % of liquids in water varies between wide limits. Some, such as alcohol and glycerine, are soluble in all proportions. Oils, such as kerosene, are practically insoluble; hence the old adage, " Oil and water will not mix." Carbon disulphide is also insoluble, as may be seen by the formation, after agitation, of two distinct layers of liquid. The existence of two layers, however, is not always absolute proof of insolubility. JEther and water form two layers, but each dissolves appre- ciably in the other. In many cases a rise of temperature increases the solubility of liquids in water. Solutions of Solids. The solubility of solids in water is a matter of tremendous practical importance. The abundance of water and its power to dissolve such a vast number of different solids have led some to call water " the universal solvent." The far-reaching effect of this marvelous power in nature and its indispensable value to man have been considered. (See above.) The degree of solubility of solids in water varies with the substance and with the temperature of the water. Some, like potassium permanganate, are very soluble, while others, like calcium sulphate, are difficultly soluble. In most cases solubility increases with a rise of temperature ; hence the common practice of heating to hasten solution. The effect of increased temperature on solubility is some- times very marked, the solubility being increased fourfold in some cases. Calcium hydroxide is less soluble in hot than in cold water, while common salt (sodium chloride) dissolves to about the same degree in each. There is a limit to solubility. That is, a given volume of water at a fixed temperature will dissolve a definite weight of solid and no more, although some undissolved solid remains in the water. 44 Descriptive Chemistry. TABLE OF SOLUBILITY OF SOLIDS IN WATER. NUMBER OF GRAMS SOLUBLE IN 100 GRAMS SOLIDS. OF WATER AT 20 C. IOOC. Calcium chloride 74 155 Copper sulphate (cryst.) . . 42.3 203.3 Magnesium sulphate . . 36.2 73.8 Potassium chlorate .... 7.2 59-5 Potassium chloride .... 35 57 Potassium dichromate . . . 13 102 Potassium nitrate . . . 3'-7 2 4 6 Potassium sulphate .... 10.6 26 Sodium chloride 36 39-7 A solution is saturated at a given temperature when it will dissolve no more solid. If a hot solution, especially one which contains much solid, is cooled slowly, the solid soon begins to separate from the liquid, since solubility usually decreases with a fall of temperature. Often the solid is deposited in masses having a definite shape. This operation is called crystallization, and the masses are called crystals (see below). The shape and color of the crystal are characteristic of the substance, and serve to identify it. Thus, common salt crystallizes in cubes. Sometimes it is more convenient to evaporate a hot, con- centrated solution. The point of saturation at the lower temperature is thus reached so gradually that the crystals can grow symmetrically. A brief account of crystals will be found in 3 of the Appendix. A solid can also be separated from a solution by precipi- tation. This may be done in two ways, (i) By adding a liquid in which the solid is not very soluble. Thus, when Properties of Water. 45 water is added to an alcoholic solution of camphor, the liquid becomes turbid, or cloudy, because the camphor is not soluble in water. That is, the solid has been precipi- tated as very fine particles which remain suspended in the liquid for some time. Since the separated solid sooner or later falls to the bottom of the vessel, it is called a precipi- tate. (2) By changing the dissolved solid into another substance not soluble in the liquid. Such chemical changes are examples of double decomposition. Thus, when so- dium chloride solution is added to silver nitrate solution a white, curdy precipitate of silver chloride is formed. A soluble silver compound has thus been changed into an insoluble silver compound, thereby removing the combined silver from the solution. So, also, a soluble chlorine com- pound (sodium chloride) has been changed into an insoluble chlorine compound (silver chloride), thereby removing the combined chlorine from the solution. Precipitation is a very common operation in chemistry. A hot, saturated solution of some solids, such as sodium sulphate and sodium thiosulphate, deposits no crystals when the clear solution cools. Such solutions are super- saturated. Supersaturation can occur only when the un- dissolved solid is not present. Hence, if a fragment of the solid is dropped into the supersaturated solution, crys- tals soon begin to form upon the fragment, and this sepa- ration continues until nearly all the substance is deposited, often forming a solid mass in the test tube. Dust, or even shaking, causes the substance to be deposited, hence the solution should be kept corked and left undisturbed. Sat- uration is analogous to stable equilibrium, while supersatu- ration resembles unstable equilibrium. Water of Crystallization. Crystals deposited from the water solution of many solids, even after they are dried 4.6 Descriptive Chemistry. by pressing between filter paper or by exposure to a mod- erate temperature, often contain water which seems to be an essential part of the chemical compound. This water is called water of crystallization. The crystals of some compounds, e.g. sodium carbonate and sodium sulphate, lose their water of crystallization and crumble on exposure to the air. This property is called efflorescence, and such crystals are said to effloresce or to be efflorescent. Heat will drive the water of crystallization from crystals which contain it, e.g. gypsum, alum, and copper sulphate. The proportion of water of crystallization in crystals is not arbitrary. It is constant in the same compound when crystallized under uniform conditions, but the proportion varies between wide limits in different substances. No explanation has been given of the varying amount of water of crystallization, nor of its necessity for the form and color of some crystals and not for others. Some well-crystallized substances contain no water of crystallization, e.g. potassium nitrate, potassium dichromate, sugar, and salt. Crystals which have lost their water of crystallization are said to be dehydrated or anhydrous. Thus, the grayish powder obtained by heating the blue crystallized copper sulphate is called dehydrated cop- per sulphate. The words dehydrated and anhydrous have been extended to describe any substance from which water has been removed, as anhy- drous alcohol or ether. The opposite term, hydrated, is sometimes applied to a compound to emphasize the fact that it contains water of crystallization. Deliquescence. Many substances, crystallized and uncrystallized, absorb water when exposed to the air, and become moist, or even dissolve in the water. Calcium chloride, potassium carbonate, zinc chloride, sodium hydrox- ide, and potassium hydroxide belong to this class. This property is called deliquescence, and the substances are said to deliquesce, or to be deliquescent. The term hygro- scopic is applied to substances which absorb water, but hygroscopic substances do not dissolve in the absorbed Properties of Water. 47 water, and sometimes do not even become moist. Quick- lime is hygroscopic. Common salt, or sodium chloride, often appears to deliquesce, espe- cially in damp weather. The deliquescence is due, however, to the presence of magnesium and calcium chlorides. Sodium nitrate is some- what deliquescent, and cannot be used in the manufacture of gunpowder, so potassium nitrate is used instead. This property of deliquescence is often utilized in the laboratory to remove water vapor from gases, cal- cium chloride being especially serviceable for this purpose. Thermal Phenomena of Solution. Solution is often accompanied by an appreciable change of temperature. When sulphuric acid is poured into water, heat is produced. With large quantities the heat is so great that the mixture often boils, and sometimes the hot acid is spattered. Hence, the acid should be added slowly to the water, and the mixture constantly stirred. Other substances which dissolve with the liberation of heat are fused calcium chloride, potassium hydroxide, and sodium hydroxide. Some which dissolve with a fall of temperature are crystal- lized calcium chloride, ammonium nitrate, ammonium chloride, and potassium nitrate. This subject is still under investigation. Solution and Chemical Action. Probably when a sub- stance dissolves it is so modified that it can participate more readily in chemical changes. Hence, solution is an aid to chemical change, and is often an easy means of causing it. Thus, if dry tartaric acid and sodium bicar- bonate are mixed, there is no evidence of chemical action ; but when the mixture is poured into water, the copious evolution of carbon dioxide gas is conclusive evidence of a chemical change. Similarly, when a dry mixture of ferrous sulphate and potassium ferrocyanide is poured into water, the immediate appearance of a blue precipitate shows that the water was needed for the chemical change. Solution is such an important aid to chemical action that many substances employed in the laboratory are in solution, and many processes in chemistry are " wet " processes. 48 Descriptive Chemistry. Mention has already been made of the application of this fact to many industries. The Nature of Solution has long been a subject of specu- lation and study. The problem as a whole is still unsolved, though much light has been thrown upon the question by recent investigations (see Chapter X). EXERCISES. 1 . Mention several familiar properties of water. 2. In what forms does water exist ? 3. Give the per cent of water in some familiar foods. 4. Develop the topics : () water is an erosive agent ; () water is a solvent in nature ; (V) water has many industrial applications ; (d) water behaves exceptionally when heated from o C. to ioC.; (e) ice floats ; (/) water is a cleansing agent. 5. Explain these expressions : (a) water has its maximum density at 4 C. ; (b) the density of ice is 0.92 ; (V) steam is invisible ; (d) the lower the pressure, the lower the boiling point ; (e) 10 cc. per liter ; (_/") parts per million. 6. How do natural waters illustrate the solvent power of water? 7. What is (a) mineral water, () soft water, (c) hard water, (d) sulphur water, (e} chalybeate water ? 8. What does ocean water contain? Why is the sea water salt? 9. What constitutes a safe drinking water? How may city water be purified ? What is a water analysis ? 10. Describe the operation of distillation. What is a condenser and why is it" so named? Is distillation a new or an old process? Of what industrial use is it? 11. Define and illustrate (a) water of crystallization, (b) efflores- cence, (c} deliquescence, (d) hygroscopic, (e) anhydrous, (/) dehy- drated, (g) crystal, (//) crystallization. 12. Define and illustrate (a} solution, (b} solvent, (V) solute, (d) soluble, (e) slightly soluble, (/) very soluble, (g} insoluble, (h} di- lute, (/) concentrated, (/) saturated solution, () supersaturated solution. 13. Give several facts about the solubility of gases in water. What is (a) soda water, (b} carbonated water? How do we know that air Properties of Water. 49 dissolves in water? Why do subterranean waters often contain dis- solved gases? State Henry's law of the solubility of gases. What effect has (a) heat and () cold on the solubility of gases in water? 14. What liquids are soluble in water? How may such liquids be separated from water? 15. What general effect has (a} heat and () cold on the solubility of solids in water? Mention some solids which are (a) very soluble, (b) moderately soluble, (<:) almost insoluble in water. Develop the topic : There is a limit to the solubility of solids in water. 1 6. (a} How would you find the approximate amount of water in (i) milk, (2) an apple? (b) How would you find the per cent of each substance in a mixture of sand and sugar? 17. Develop the topic: Solution aids chemical change. Why are so many solutions used in a laboratory ? 1 8. What changes in volume occur when (a} ice melts, () water freezes, (c) water is heated from oC. to i5C., (d) water is cooled from I5C. to oC.? 19. Write an essay on "Mineral Springs in the United States." PROBLEMS. 1. If 1.5 gm. of crystallized barium chloride lose 0.22 gm. when heated to constant weight, what per cent of water of crystallization does it contain? 2. If 2 gm. of another lot of barium chloride lose 0.295 gm., what per cent of it was water of crystallization ? 3. If a liter of sea water has a density of 1.25, how many grams of "salt" does it contain? 4. If the density of ice is 0.92, what volume will a liter of water at 4C. occupy when frozen? Ans. 1.087 ! 5. How much water (approximately) is contained in (a) 2 Ib. of lobster, (b} 56 Ib. of potatoes, (c) i Ib. of tomatoes, {d} 2 Ib. of milk, (e) i Ib. of white bread, (/) a human body weighing 150 Ib. ? 6. If a kilogram of sea water contains 36.4 gm. of "salt," what per cent of the water is " salt " ? 7. If a block of ice weighs 280 kg., what is its volume? Ans. 304.3 1. 8. A solution measures 100 cc. and contains 15 gm. of potassium nitrate. What per cent of water and of solid is in the solution ? CHAPTER V. COMPOSITION OF WATER. - WATER was considered an element until about the end of the eighteenth century. At that time it was shown to be a compound of hydrogen and oxygen. Many famous chemists worked on this problem. The Composition of a Compound is determined either by analysis or synthesis, i.e. by taking it apart or putting its parts together. Sometimes both methods are used, since each method fortifies the other and strengthens the final conclusion. These methods find excellent application in determining the composition of water. Analysis and synthesis may be qualitative or quantitative. A quali- tative experiment is a study of the properties of elements and com- pounds with a view of discovering what they contain. A quantitative experiment is an accurate determination of the weight or volume of the components of a compound. Qualitative tests involve merely quality, while in quantitative tests quantity is the essential feature. Obviously, a complete determination of the composition of a compound requires both tests. Water contains Hydrogen. When steam is passed over heated metals, hydrogen is liberated. Lavoisier's demonstration of this fact has already been considered (see Preparation of Hydrogen). The fact that sodium liberates hydrogen from water at the ordinary temperature has also been discussed (see ibid.). If red litmus paper is put into the water from which the sodium has liberated hydrogen, the litmus paper becomes blue. This change 50 Composition of Water. 51 i of color from red to blue shows that an alkali is in the water, because alkalies turn red litmus paper blue. The alkali is sodium hydroxide, and it may be obtained as a white solid by evaporating the water. Sodium hydroxide is a compound of sodium, hydrogen, and oxygen, and is formed by replacing part of the hydrogen of water by sodium. Since sodium liberates hydrogen from water, and forms at the same time a compound sodium hydroxide containing hydrogen, the hydrogen in water must be divisible into two parts. Now if o. I gm. of sodium is allowed to act upon water, 48.22 cc. of hydrogen are liber- ated ; and if the sodiunThydroxide thus formed is dried and heated with sodium, 48.22 cc. more of hydrogen are ob- tained. This shows that the hydrogen in water is divisible into two equal parts a fact which will soon be utilized. Water contains Oxygen. The fact that oxygen is a component of water has already been suggested, e.g. (i) by the production of water when hydrogen is burned in air, (2) by the formation of a compound of iron and oxy- gen when steam is passed over hot iron, and (3) by the formation of sodium hydroxide when sodium acts upon water. These proofs, however, are all indirect. A simple direct demonstration of the presence of oxygen in water may be made by allowing chlorine water to stand in the sunlight. (Chlorine water is prepared by saturating water with chlorine gas an element to be studied in Chapter XI.) A long tube like that shown ^ in Figure 9 is completely filled GU with chlorine water, the open end is FIG. 9. Tube for decompo- , L . . sition of water by chlorine. immersed in a vessel containing some of the* same solution, and the whole apparatus is placed in the direct sunlight. Bubbles of gas soon appear in the liquid, and after a few hours a small volume of Descriptive Chemistry. gas collects at the top of the tube. This gas may be shown, by the usual tests, to be oxygen. The Electrolysis of Water is its decomposition by elec- tricity. It is accomplished in the apparatus shown in Figure i o. Since pure water does not conduct electricity, sulphuric acid is added. Enough of this acid mixture is poured into the apparatus to fill the reservoir half full after the stopcocks have been closed. As soon as an electric battery of two or more cells is connected by wires with the piece of platinum near the bottom of each tube, bubbles of gas form on the platinum, and as the action proceeds, the bubbles rise and displace the water in each tube. The volume of gas is greater in one tube. Assuming that the tubes have the same diameter, the volumes are in the same ratio as their heights, which will be found by measurement to be two to one. The larger volume of gas is FIG. io. Hofmann apparatus for hydrogen and the smaller one electrolysis of water. is oxygen. Many accurate repe- titions of this experiment have shown that only hydrogen and oxygen are produced, and that the ratio of their volumes is two to one. It has also been shown that the sum of the weights of the two gases equals the weight of the water decomposed. The whole experiment demonstrates that Composition of Water. 53 water is a compound consisting of two volumes of hydro- gen combined with one volume of oxygen. Water was first decomposed by electricity in 1800 by Nicholson and Carlisle. Davy confirmed their work by a series of brilliant experi- ments extending through a period of six years (1800-1806). During this time he not only proved that the volume of hydrogen is double that of oxygen, but by electrolyzing water in a gold vessel placed in an atmos- phere of hydrogen, he proved that nothing but these gases is produced. The Quantitative Composition of Water. The fore- going facts about the composition of water have been mainly qualitative. They have shown by analysis and synthesis that water consists of hydrogen and oxygen, and that the ratio of their volumes is approximately two to one. Decisive evidence of the quantitative composition of water is obtained by a determination of its volumetric and its gravimetric composition. Volumetric means "by volume" and gravimetric means " by weight." The Volumetric Composition of Water is determined by exploding a mixture of known volumes of hydrogen and oxygen in a eudiometer. Gas volumes which are to be compared with each other must be dry and at the same temperature and pressure. This requirement, which is called the " standard condition," is inconvenient, and almost impracti- cable. Hence, it is customary to measure each volume of moist gas under the existing conditions, and then reduce the observed volume to that volume which the gas would occupy if standard conditions pre- vailed. The reduction to standard conditions is accomplished by the formula j/r /pi _ n \ 760(1 + . 00366 /) In the formula l V = the corrected volume. V = the observed volume. 1 A complete discussion of the laws of gases, the principles which control their measurement, together with the development of the above formula for reduction to standard conditions, may be found in Appendix B of the author's " Experimental Chemistry." See also the Laws of Boyle and Charles in Chapter II, and Vapor Density in Chapter IV (this book). 54 Descriptive Chemistry. P f = the observed pressure. / = the observed temperature. a the vapor tension at / C. A convenient form of apparatus for determining the volu- metric composition of water is shown in Figure n. The essential part is the eudiometer, F. In this graduated glass tube the gases are accurately measured and ex- ploded. The electric spark which causes the explosion is obtained from an induc- tion coil and battery. The spark leaps across the space between the platinum wires at the top of the eudiometer, and the heat produced by this spark causes the hydrogen and oxygen to combine and form water. Oxygen and hy- drogen are introduced separately into the eudi- ometer, measured, and - exploded. After the FIG. ii. Apparatus for determining the volu- explosion, which IS indi- cated by a slight click or flash of light, water from the reservoir, E, rushes up into the eudiometer. The water does not completely fill the tube, because an excess of one gas is added. This additional gas takes no part in the chemical change, but merely serves to lessen the violence of the explosion, which otherwise might break the eudiometer. The quantity of water formed by the union of the hydrogen and oxygen Composition of Water. 55 is too minute to measure. Repeated trials of this experi- ment show that two volumes of hydrogen always combine with one volume of oxygen. This is the volumetric com- position of water. The discovery of the volumetric composition of water was not made by one chemist alone. Priestley, about 1780, noticed that when a mixture of air and hydrogen was exploded, " the inside of the glass, though clear and dry before, immediately became dewy." Cavendish, in 1781, showed that when a mixture of two parts hydrogen and one part oxygen was exploded, nothing but water was formed. Watt, in 1783, was the first to state that water is a compound, though he per- formed no experiments and probably did not understand the real nature of its components. Lavoisier in the same year verified many facts pre- viously noticed but not completely understood, and undoubtedly first clearly recognized and stated what his contemporaries had overlooked. The final proof of the volumetric composition of water was an accurate verification in 1805 by Gay-Lussac and Humboldt of the previous ob- servation that two volumes of hydrogen unite with one volume of oxygen. The Gravimetric Composition of Water is determined by passing dry hydrogen over copper oxide. The method depends upon the fact that many oxides, such as those of lead, copper, and iron, when heated in a current of hydro- =7) (e= = l i i c c' B FIG. 12. Apparatus for determining the gravimetric composition of water. gen, give up their oxygen, or, chemically speaking, these oxides are reduced to metals. By this reduction the oxy- gen of the oxide combines with the hydrogen, thereby forming water which is collected in a weighed tube. 56 Descriptive Chemistry. A convenient form of apparatus is shown in Figure 12. The copper oxide is placed in the combustion tube, CC, which is made of hard glass. The Marchand tube, D, which is filled with calcium chloride, collects and retains the water formed in the combustion tube, as the hydrogen passes over the hot copper oxide. The tubes A, B, and E keep moisture out of the apparatus. The experiment is very simple. Copper oxide is placed in the combustion tube, which is then carefully weighed. The Marchand tube, being filled with calcium chloride, is also weighed. After the other tubes are properly filled and the hydrogen generator adjusted, the tubes are connected as shown in the figure. The combustion tube is now heated, and mois- ture collects in it; as the heat increases the copper oxide glows, and the moisture passes into the Marchand tube. When the operation is over and the apparatus is cool and free from hydrogen, the combustion tube and Marchand tube are weighed. The gain in weight of the Marchand tube is the weight of the water formed, while the loss in weight of the combustion tube is the weight of the oxygen contained in this water. An illustration will make this clear. Dumas and Stas, who performed this experiment accurately in 1843, found substantially that the combus- tion tube lost 5.251 gm. of oxygen, while the Marchand tube gained 5.909 gm. of water. But 5.251 and 5.909 are in the same ratio as 8 and 9. Thus : 5.251 : 5.909 : : 8 : 9. This means that oxygen makes up f of water. The re- maining ^ is of course hydrogen. In other words, the gravimetric composition of water is eight parts oxygen and one part hydrogen. This ratio is often stated in per- centage ; thus water contains Composition of Water. 57 88.88 per cent of oxygen. 1 1 . 1 1 per cent of hydrogen. For reasons which will soon be given, it is more conven- ient to state the composition of water by weight, as two parts hydrogen to sixteen parts oxygen. The gravimetric composition of water was first determined about 1820 by Berzelius and Dulong. Their work was verified by Dumas and Stas in 1843. A Comparison of the Volumetric and Gravimetric Com- position of Water shows that the results of the two methods agree. One volume of oxygen is sixteen times heavier than an equal volume of hydrogen (see Density of Hydrogen). Therefore, the one volume of oxygen must be eight times heavier than the two volumes of hydrogen in water. That is, the oxygen in water weighs eight times more than the hydrogen. But this is the ratio actually found in determining the gravimetric composition of water by an independent experiment. These facts strengthen our belief that the composition of water is By weight, one part hydrogen and eight parts oxygen. By volume, two parts hydrogen and one part oxygen. Summary. The following facts have been shown con- cerning the composition of water : (1) Water is a chemical compound 'of hydrogen and oxygen. (2) It is formed when hydrogen is burned in air, or when a mixture of hydrogen and oxygen is exploded. (3) It can be decomposed by electricity into hydrogen and oxygen in the ratio of two volumes of hydrogen to one volume of oxygen. 58 Descriptive Chemistry. (4) Sodium liberates hydrogen from water and forms at the same time a solid containing a quantity of hydrogen equal to the quantity of hydrogen liberated. Iron, other metals, and carbon liberate hydrogen from water, forming at the same time an oxide of the respective substance. (5) Chlorine liberates oxygen from water. (6) Two volumes of hydrogen, when exploded with one volume of oxygen, combine to form water, and the weight of the water formed equals the weight of the gases used. (7) Water is formed by the union of two parts by weight of hydrogen and sixteen parts by weight of oxygen. EXERCISES. 1. How is the composition of a compound determined ? 2. Define (a) synthesis, () analysis, (<:) qualitative, (//) quantita- tive, (e) volumetric, (/) gravimetric. 3. How would you prove that water is composed of hydrogen and oxygen ? 4. How do we know that the hydrogen in water is divisible into two equal parts ? 5. What is the electrolysis of water ? How is it accomplished ? What does it prove about the composition of water ? When and by whom was it first performed ? What did Davy contribute toward the solution of the problem ? 6. What is the volumetric composition of water ? How is it deter- mined ? Who worked on this problem, and what did each contribute to its solution ? 7. Answer the same questions (as in 6) about the gravimetric com- position of water. 8. Compare the volumetric and the gravimetric composition of water. 9. What does the burning of hydrogen show about the composition of water ? 10. Summarize the essential facts regarding the composition of water. Composition of Water. 59 ii. Give a brief biographical account of (a) Nicholson and Carlisle, () Dumas, (c} Humboldt, (d} Stas, (e) Watt, (/) Gay-Lussac (see Appendix, 4) . PROBLEMS. 1. What weight of (a) hydrogen and (fr) oxygen can be obtained by decomposing 125 gm. of water ? 2. What volume of (a) hydrogen and (6) oxygen can be obtained by decomposing 9 1. of water ? 3. What weight of hydrogen must unite with 16 gm. of oxygen to form water ? What weight with (#) 40 gm., (b) 70 gm., (c) 160 gm. ? 4. What volume of oxygen must unite with 2 1. of hydrogen to form water ? What volume with (a) 40 1., () 40 cc., (c) 40 qt, (d ) 95 vol- umes, (e) 1 60 1. ? 5. What volume of oxygen is necessary to unite with 100 gm. of hydrogen to form water ? (Suggestion : What is the weight of a liter of oxygen ?) 6. Hydrogen is passed over 2.48 gm. of hot copper oxide, which at the end of the experiment weighed 2.24 gm. ; the water formed weighed 0.27 gm. In what ratio did the hydrogen and oxygen combine ? 7. Berzelius and Dulong, in 1820, obtained the following results in their determinations of the gravimetric composition of water : Loss of weight of copper oxide (in grams), 10.832 and 8.246. Weight of water formed, 12.197 and 9.27. Calculate in each case the ratio in which the hydrogen and oxygen combined. What is the average ratio ? 8. Dumas and Stas repeated the above work in 1843, and found as an average of nineteen determinations, that 840.161 gm. of oxygen formed 945.439 gm. of water. Calculate the ratio of combination. Hydrogen Dioxide is a liquid composed of hydrogen and oxygen. But the proportion of the components is not the same as in water. It contains two parts of hydrogen and thirty-two parts of oxygen by weight. It is often called, especially in commerce, hydrogen peroxide, because its relative proportion of oxygen is greater than in water the other hydrogen oxide. It is manufactured by treating barium dioxide (or peroxide) with sulphuric or hydrochloric acid. The commercial solution has a vari- able strength, and usually contains three or more per cent of hydrogen dioxide. It has a sharp, pungent odor, and a bitter, metallic taste. 60 Descriptive Chemistry. Hydrogen dioxide is an unstable compound ; it decomposes slowly at the ordinary temperature, and very rapidly if heated. The dilute, com- mercial solution is somewhat stable, but heat decomposes it completely into water and oxygen. The ease with which it yields oxygen makes it a good oxidizing agent. In this respect, hydrogen dioxide resembles ozone, and, indeed, they are sometimes mistaken for each other. It is also a reducing agent, and is frequently used as such in the laboratory. It is used extensively to bleach animal and vegetable matter, such as human hair, ostrich feathers, fur, silk, wool, cotton, bone, and ivory. It is also used as an antiseptic and disinfectant in surgery. Large quanti- ties are used to restore the color to faded paintings a use suggested by The'nard, the discoverer. In the laboratory it is proving a service- able reagent. Hydrogen dioxide is found in the air, in rain and snow, but the proportion is variable and exceedingly small. CHAPTER VI. THE ATMOSPHERE NITROGEN. The Atmosphere is the great mass of gas surrounding the earth and extending into space. Its estimated height is fifty to several hundred miles. We live at the bottom of this vast ocean of air, as it is often called. Aristotle (384-322 B.C.) regarded air as one of the four elementary principles whose combinations made up all substances in the universe. The other three were earth, fire, and water. He taught that air pos- sesses two fundamental properties, heat and dampness. The early chemists used the word air in the sense in which the word gas is now employed. Thus, we have already learned that hydrogen was first called inflammable air. The terms atmosphere and air are often used inter- changeably, though by air we usually mean a limited por- tion of the atmosphere. Many skillful chemists have studied the action of air on living things, its relation to combustion, the effect of its weight, its composition, and its varied properties. Their work has contributed many fundamental facts to science. General Properties of the Atmosphere. Air has weight. We often use the expression " light as air." But a cubic foot of air weighs 1.28 oz. and a room 40 x 50 X 25 ft. contains about two tons of air. The total weight of the atmosphere has been estimated to be five thousand millions of millions of tons. This enormous mass resting upon the earth exerts a pressure which is about fifteen pounds on every square inch. This amount of pressure upon a 61 62 Descriptive Chemistry. square inch is called "an atmosphere," and it is some- times used as a unit of pressure. Thus, three atmospheres means a pressure of forty-five pounds per square inch. It is this pressure which causes water to rise in pumps and flow through siphons. Atmospheric pressure is exerted in all directions and is variable. It is measured by the barometer. The normal or standard pressure of the at- mosphere is equal to the weight of a column of mercury one square inch in cross section and 29.92 in. high, or one square centimeter in cross section and 760 mm. high. But since atmospheric pressure is at the rate of fifteen pounds to the square inch, it is necessary to know the height only of the mercury column in order to know the pressure. The pressure of the atmosphere varies as the height and the compo- sition of the atmosphere vary, and the barometer changes accordingly. The weight of a liter of dry air at o and 760 mm. is i .293 gm. The appreciable movements of the atmosphere are the winds. Ingredients of the Atmosphere. The atmosphere is a mixture of several gases. But since this mixture always contains about 78 parts of nitrogen and 21 parts of oxygen by volume, we often speak of air as consisting solely of these two gases. Besides this large proportion of oxygen and nitrogen, the air always contains small and variable proportions of water vapor and carbon dioxide gas. Be- sides these four ingredients, air always contains the gases argon and helium, and usually ozone, hydrogen, hydrogen peroxide, compounds related to ammonia and nitric acid, dust, and germs. The composition varies but slightly in different localities. Near the city air may contain a rela- tively larger proportion of dust, ammonia, sulphur com- pounds, and acids ; in the country the proportion of ozone is relatively large ; at the ocean the air contains consider- able salt. The Atmosphere Nitrogen. 63 General Properties of Nitrogen. The chemical ele- ment, nitrogen, constitutes about 78 per cent of the atmos- phere (by volume). It is a colorless gas, and has no taste or odor. It is somewhat lighter than air, and is very slightly soluble in water. In many respects it differs markedly from oxygen. Thus it will not support combus- tion, neither will it burn nor sustain life. Animals die if left in nitrogen. * The fact that a candle flame quickly goes out and a mouse soon dies in nitrogen was first observed by Rutherford, an English physician, who discovered the gas in 1772. Soon after, Lavoisier showed the true relation of nitrogen to the atmosphere. To emphasize the inability of the gas to support life, he called the new gas azote, the name now used for it by some French chemists. Nitrogen is not poisonous, for a large proportion of the air we breathe is nitrogen. Its function in the atmosphere is to dilute the oxygen. It is an inert element. It com- bines with only a few other elements, and many of its compounds easily decompose. Oxygen and Nitrogen in the Atmosphere. The chem- ical activity of the atmosphere is due to the free oxygen it contains. We have already learned that oxygen is an i active chemical element. If the air were largely oxygen, rusting and decay would proceed with astounding rapidity, and fires once started would burn with .great violence. On kthe other hand, nitrogen is inactive. And if the air con- tained much more than the normal amount, chemical action would be slower. Oxygen alone is too active, while nitrogen alone is inactive. To be serviceable to man, oxygen must be diluted with nitrogen, while nitro- gen must be accompanied by a small proportion of oxygen. 64 Descriptive Chemistry. The presence of oxygen and nitrogen in the atmosphere, and the functions of the two gases, were first clearly explained by Lavoisier in 1 777, though many others Boyle, Priestley, Rutherford, and Scheele helped solve the problem. Composition of the Atmosphere. Samples of air from various parts of the globe show a remarkable uniformity of composition. Until 1895 it was supposed that pure air consisted solely of oxygen and nitrogen. But it has been found that about one per cent of the gas hitherto called nitrogen is argon, a gas so much like nitrogen, and so difficult to separate from the latter, that for years it had been overlooked (see Argon, below). According to the most recent results, the following is THE COMPOSITION OF PURE DRY AIR. INGREDIENT. PERCENTAGE. By volume. By weight. Nitrogen 78.06 21.00 0.94 7#Y -23.2 .V? Oxvefen Argon .... \ The composition of the atmosphere was studied by Priestley, but his results were conflicting. Cavendish, in 1781, was the first to show that the proportion of oxygen and nitrogen in air is nearly constant. Since his time this result has been confirmed by many chemists, especially by Bunsen, who is widely known as the inventor of the Bunsen burner, which is used as a source of heat in chemical laboratories. The Volumetric Composition of the Air may be found by introducing a known volume of pure air into a eudiom- eter and exploding it with a known volume of hydrogen. The oxygen of the air combines with twice its volume of hydrogen, forming a minute quantity of water ; hence one The Atmosphere Nitrogen. third of the diminution in volume is the volume of oxygen in the air. The difference between the volume of oxygen found and the original volume of air is the volume of nitrogen. An illustration will make this experiment clear. Suppose (i) we mix and explode loocc. of air and 50 cc. of hydrogen, or 15000. in all, and (2) that the residue measures 87 cc. Now, 150 87 = 63, hence 63 cc. of the total volume combined to form water. But one third of 63 cc. is oxygen, which came from the original volume of air. Hence, 63 -r- 3 = 21, the volume of oxygen in 100 cc. of air. The remainder, 79 cc., is nitrogen, argon, and other gases. Another Method, <;often used to determine the volu- metric composition of the air, is based on the fact that phosphorus will com- bine slowly with oxygen, even at the ordinary tem- perature. The operation is performed in an apparatus like that shown in Figure 1 3. A piece of phosphorus, C, attached to a wire, is inserted into a graduated glass tube, />, containing a measured volume of air. White fumes indicate im- mediate action. These fumes are solid particles of an oxide of phosphorus FIG. 13. Apparatus for determining the com- position of air by phosphorus. called phosphorus pentoxide. 'They soon dissolve in the water, which rises higher in the tube, as the oxygen combines with the phosphorus. In a few hours the phosphorus is removed, and the volume of gas is read. The difference between the first and last volumes is oxygen. The gas remaining in the tube is, of course, a mixture of nitrogen and argon. In performing this experiment unusual care must be taken not to touch the phosphorus with the bare hands. 66 Descriptive Chemistry. The Gravimetric Composition of Air was first accurately determined in 1841 by the French- chemists, Dumas and Boussingault. The average result of many experiments tTT'O o Lm Oxygen . . . . 23 parts by weight. Nitrogen ... 77 parts by weight. We know, however, that the correct proportions are Oxygen .... 23.2 parts by weight. Nitrogen . . . 75.5 parts by weight. Argon .... 1.3 parts by weight. They passed pure air through a weighed tube containing copper, and arranged so that heat could be applied. The oxygen of the air com- bined with the copper, while the nitrogen passed on into a weighed globe. Both tube and globe increased in weight. The increase in the tube was the weight of the oxygen, while the increase in the globe was the weight of the nitrogen. Water Vapor in the Atmosphere. Water vapor is always present in the atmosphere, owing to the constant evaporation from the ocean and other bodies of water. The total amount present is large, though variable. A given volume of air will absorb a definite volume of water vapor and no more, and the amount depends largely upon the temperature. Air containing its maximum amount of water vapor is said to be saturated at that temperature, or to contain 100 per cent of water vapor. The saturation point is also called the dew point. On a pleasant day the relative humidity of the air, i.e. the amount of water vapor present, may vary from 30 to 90 per cent, the aver- age being about 50 per cent. Warm air holds more vapor than cool air. The amount of water vapor in the air has a marked influence on the physical condition of man. The depressing weather during " dog days " is due to the The Atmosphere Nitrogen. 67 high relative humidity of the air, which sometimes reaches 95 per cent. The absence of life in deserts is largely due * to the dry air .above them. Much of the languor felt in a " close " room or crowded hall is partly caused by the excess of water vapor in the "bad" air. The presence of water vapor in the air is shown by the moisture which col- lects on the outside of a vessel containing cold water, such as a pitcher of iced water. The moisture comes from the air around the vessel. For a similar reason, water pipes in a cellar and the cellar walls themselves are moist in summer. The deliquescence of calcium chloride, common salt, and other substances likewise reveals the presence of water vapor in the air (see Deliquescence). When the temperature of the air falls, the water vapor condenses and is deposited in the form of dew, rain, fog, mist, frost, snow, sleet, or hail. The clouds are masses of water vapor which has been condensed by the cold upper air. Carbon Dioxide in the Atmosphere. Carbon dioxide is one product of the respiration of animals, and of the combustion and decay of organic substances. By these processes an immense quantity of carbon dioxide is being constantly poured into the atmosphere. The quantity in the atmosphere is variable, though not between such wide limits as the water vapor. The proportion in normal air is about 4 parts in 10,000 parts of air. Over the ocean the proportion is smaller, but in the air of cities it is greater. In crowded rooms the proportion is often as high as 33 parts in 10,000, because carbon dioxide is exhaled faster than it can be removed. The proportion of carbon dioxide in the atmosphere as a whole is practically constant, largely owing to the fact that this gas is an essential food of plants (see Carbon Dioxide). The pres- ence of carbon dioxide in the air is detected by limewater. 68 Descriptive Chemistry. If Hmewater is exposed to the air, the carbon dioxide unites with the lime in the limewater, forming a thin, white crust of insoluble calcium carbonate on the surface of the Hmewater. If air is drawn through lime- water, the liquid becomes milky, because the particles of calcium carbon- ate are suspended in the liquid. The purity of air is often determined by finding out what proportion of carbon dff>xide it contains. If a known volume of dry air is drawn through a known weight of Hmewater or similar liquid, the increase in weight will be the weight of carbon dioxide in the volume of air used. The different gases in the atmosphere are not arranged in layers according to their densities. They are in con- stant circulation (see Diffusion). Hence carbon dioxide, though heavier than oxygen and nitrogen (volume for vol- ume), does not remain nearest the ground, but is distrib- uted through the air. In a few exceptional localities, carbon dioxide arises from volcanoes faster than it can diffuse, and fills the adjacent valley. Argon in the Atmosphere. Argon is a colorless, odor- less gas. Its chief characteristic is its chemical inactivity. No compounds of argon have as yet been prepared or discovered. The name argon is happily chosen, being derived from Greek words signifying inert. It constitutes 0.94 per cent by volume of the atmosphere, or 1.3 per cent by weight. Argon was discovered in 1894 by Rayleigh and Ramsay. Rayleigh had found that nitrogen from air weighed more than an equal volume of nitrogen obtained from compounds of nitrogen. Consequently, they believed that the nitrogen from air contained another gas hitherto over- looked. A series of elaborate experiments showed that after all the oxygen and nitrogen was removed from purified air, there still remained a small quantity of a new gas, which they called argon. It may be pre- pared (i) by passing pure air over healed copper to remove the oxygen, and then the remaining gas over heated magnesium or calcium to remove the nitrogen ; or (2) by passing electric sparks through a mixture of air and oxygen, and removing the compound of oxygen and nitrogen as fast The Atmosphere Nitrogen. 69 as it is formed. The latter method is a repetition of the one used by Cavendish when he determined the composition of air, and he would have no doubt discovered argon had he continued his investigations. Inert Gases in the Atmosphere. Helium, neon, krypton, and xenon have recently been discovered by Ramsay. At present little is known about these gases. They resemble argon in being inactive chemical elements. They constitute an exceedingly minute proportion of the atmosphere. Helium is also found in certain rare minerals, in the gases from some mineral springs, and in the atmosphere of the sun. It is about twice as heavy as hydrogen. According to Ramsay, " it is prob- able that helium is continually escaping from the earth in small quantities in certain regions.' 1 Air is a Mixture, in spite of the fact that we speak of its "composition." Chemical compounds have two invari- able characteristics : viz., (i) their components are in a fixed proportion, and (2) their formation and decomposition are usually attended by definite evidences of chemical action, such as light, heat, change of color and form, etc. The following facts show that air is a mixture of free gases : (1) The proportion of oxygen and of nitrogen is not fixed, but varies between small limits, which may be detected by accurate analysis. (2) When nitrogen and oxygen are mixed in the propor- tions which form air, the product is exactly like air, but the act of mixing gives no evidence of chemical action. (3) When air is dissolved in water, a greater proportion of oxygen than of nitrogen dissolves. If the oxygen and nitrogen were combined in the air, the dissolved air would, of course, have the same composition as air itself. Liquid Air is a mixture of the liquefied gases which con- stituted the air used. It is a milky liquid, owing to the presence of solid carbon dioxide and ice. If these solids are removed by filtering, the filtrate has a pale blue tint. It is slightly heavier than water. It is intensely cold, its jo Descriptive Chemistry. temperature being about 200 C. It boils at about 190 C. under atmospheric pressure. If a tumbler is filled with liquid air, the latter boils vigorously, the sur- rounding air becomes intensely cold, frost gathers on the tumbler, and in a short time the liquid air will have entirely disappeared into the air of the room. If, however, the liquid air is placed in a Dewar's bulb or flask, it evaporates so slowly that some will remain in the flask several hours. The Dewar's bulb (Fig. 14) consists of two flasks, one within the other, attached at the top ; the space between the flasks is a vacuum. Sometimes the outer surface of the inner flask is coated with mercury or silver, which helps to protect the liquid air from the heat of the atmosphere. In transporting liquid air a large Dewar's bulb or similar device is FIG. 14. A Dewar's bulb. used. One form consists of a large metal can wrapped with many thicknesses of felt and slipped into a larger can covered with canvas or felt. The liquid air is put in the inner can and a loose stopper or piece of felt is placed over the mouth. The liquid may also be kept in these cans for some time with only a moderate loss, unless the surrounding temperature is exceptionally high. Liquid air, owing to its extremely low temperature, pro- duces remarkable physical changes. A tin or iron vessel which has been cooled by liquid air is so brittle that it may often be crushed with the fingers. Nearly all plastic or soft substances, including many kinds of food, when im- The Atmosphere Nitrogen. 71 mersed in liquid air, become hard and brittle, leather being the only important exception. Mercury freezes so hard in liquid air, that it may be used as a hammer to drive a nail. When liquid air is put in a teakettle standing on a block of ice, the liquid air boils vigorously. If the kettle of liquid air is placed over a lighted Bunsen burner, frost and ice collect on the bottom of the kettle, because the intense cold of the kettle solidifies the water vapor and carbon dioxide, which are the two main products of burning illuminating gas. If water is now poured into the kettle, the liquid air boils over and the water is instantly frozen ; the water is so much hotter than the liquid air that the latter boils more violently, and since its rapid evaporation causes absorption of heat, the water gives up its heat and becomes ice. Ordinary liquid air* is from one half to one fifth liquid oxygen, and will support combustion. A red-hot rod of steel or of carbon burns brilliantly in this cold liquid. Numerous applications of liquid air have been proposed, but thus far they have not passed the experimental stage. It has been suggested that it be used as a refrigerant instead of ice, for ventilating and cooling rooms, as a blasting material, for removing diseased flesh from a wound, for destroying refuse, and as a commercial source of oxygen. The last use is based primarily on the fact that as liquid air evaporates, the nitrogen passes off first, and in a short time relatively pure oxygen remains (see Oxygen). A little liquid air was produced in 1883 with considerable labor and at an enormous expense. Now it is Easily manufactured in large quan- tities at a comparatively low cost. In the older methods of preparing liquefied gases, the gas was subjected to tremendous pressure and a low temperature. At present, air is liquefied by a different method. Com- pressed air cooled by water is forced through a pipe with a small open- ing into a larger cylinder called the liquefier. As it escapes into the liquefier it expands and its temperature falls, because expansion is a cooling process. The temperature of the liquefier is thus reduced, so that the air, which continues to enter, expands at such a low temperature that it becomes a liquid. 72 Descriptive Chemistry. NITROGEN. Occurrence. Nitrogen, besides comprising four fifths of the atmosphere, is a component of nitric acid and am- monia, and of the many compounds related to them. It is also an essential constituent of animal and vegetable matter. The name nitrogen was given to the gas by Chaptal from the fact that it is a component of niter, an old name of potassium nitrate. Preparation. Nitrogen is usually obtained from the air by remov- ing the oxygen by phosphorus. A tall jar is placed over burning phosphorus contained in a shallow dish floating in a large vessel of water. The oxygen combines with the phosphorus, leaving nitrogen, more or less pure, in the jar. Other methods may be used, such as decomposing ammonium nitrite by heat, or passing air over heated copper. Additional Properties. In addition to its inertness, already men- tioned, nitrogen is a little lighter than air, and is very sparingly soluble in water. Its density is 0.972 (air = i). One liter at o C. an'd 760 mm. weighs i.256gm. One hundred liters of water dissolve only 1.5 1. at the ordinary temperature. It combines with magnesium and a few other metals at a red heat, forming nitrides. Electric sparks cause nitrogen to combine with oxygen and with hydrogen, forming ultimately nitric acid and ammonia, hence these substances or others related to them are often found in the rain which falls during a thunder storm. Relation of Nitrogen to Life. Oxygen, carbon diox- ide, and water vapor are essentially related to the life of plants and animals. Nitrogen is also vitally connected with different forms of life. Atmospheric nitrogen merely dilutes the oxygen. Although we live in an atmosphere containing such a large proportion of nitrogen, we cannot assimilate it. According to a reliable authority, " the air as it leaves the lungs contains 79.5 per cent of nitrogen," and hence cannot become a part of the body. Yet all flesh contains nitrogen, and the rejected waste products of ani- The Atmosphere Nitrogen. 73 mals are largely combined nitrogen. The nitrogen needed by animals must be in combination to become available. And it is taken in the form of nitrogenous food, such as lean meat, fish, wheat and other grains. Most plants take up combined nitrogen from the soil in the form of nitrates (compounds derived from nitric acid) or of ammonia. Hence combined nitrogen is being con- stantly taken from the soil, and in order to preserve the fertility of the soil, nitrogen must be supplied. This is done by allowing nitrogenous organic matter to decay upon the soil, or by adding to the soil a fertilizer, which is a mixture containing nitrogen compounds. Recently it has been shown that leguminous plants, such as peas, beans, and clover, take up nitrogen from the air by means of bacteria, which are in nodules on their roots. EXERCISES. 1. What is the atmosphere? What is air? What is the literal meaning of the word atmosphere? What is the wind? 2. Develop the topics: (a) atmospheric pressure, (b) occurrence of nitrogen, (c) volumetric composition of the air, (W) gravimetric com- position of the air, (e) water vapor in the atmosphere, (/") carbon dioxide in the atmosphere, (g) air is a mixture. 3. Define and illustrate the terms : (#) an atmosphere, (<) normal pressure, (c) standard pressure, (d) dew point, (e) relative humidity, (/) inert. 4. What are the two chief ingredients of the atmosphere? The per- manent ingredients ? The variable ingredients ? The ingredients found in traces? What are sometimes found in the air of cities? 5. What is the symbol of nitrogen? What are its general proper- ties? Its special properties? What is its main function in the atmos- phere? How may it be prepared? 6. When and by whom was nitrogen discovered? Why and by whom was it named "azote 11 and "nitrogen 11 ? 7. What is the relation of nitrogen to animal and to vegetable life? 74 Descriptive Chemistry. 8. Compare the functions of oxygen and nitrogen in the atmos- phere. What famous chemists helped solve this problem? 9. State the composition of pure air (a) by volume, and (b) by weight. 10. Give a brief biographical account of (a) Cavendish, (^) Dumas, (c) Rutherford. (See Appendix, 4.) 11. What is a cloud? The dew? Why does moisture gather on cellar walls? Why are mines often damp? What is (a) rain, (t>) fog, (c) mist? 12. Describe the action of air upon (a) limewater, and (b} calcium chloride. 13. How does the atmosphere illustrate the diffusion of gases? 14. What is argon? Give a brief account of (a) its discovery, () its properties, (c) its method of preparation. What proportion of pure air is argon? What is the significance of the name argon f 15. Give a brief account of helium, neon, krypton, and xenon. 1 6. What is liquid air? What are its chief properties? State briefly its method of manufacture. Describe its action (a) upon solids, such as rubber, (b) upon liquids, such as mercury, (c) upon hot steel, (d) when evaporated quickly. Describe a Dewar's bulb. PROBLEMS. 1. If a man inhales 18 cu. ft. of air an hour, what weight of oxy- gen does he consume in 24 hr. ? 2. What is the weight of air in a room, 6x6x3111., if a liter of the air weighs 1.3 gm. ? 3. A mixture of 25 cc. of air and 50 cc. of hydrogen is exploded. The residue measures 60.3 cc. What per cent of oxygen did this sample of air contain ? 4. How many kilograms of pure air are needed to yield 100 kg. of oxygen ? 5. Express in inches the following barometer readings : (a) 760 mm., (<) 740 mm., (c) 75 cm., (d) 0.749 m., (e) 7.67 dm. 6. Dumas and Boussingault, in 1841, found in a sample of air, 12 -373 g m - of nitrogen and 3.68 gm. of oxygen. What per cent of each was found? 7. What is the weight at o C. and 760 mm. of (#) 1000 cc. of dry air? Of () 750 1., (c) 1750 cc., (d) 850 cu. m.? CHAPTER VII. LAW AND THEORY LAWS OF DEFINITE AND MUL- TIPLE PROPORTIONS ATOMIC THEORY ATOMS AND MOLECULES SYMBOLS AND FORMULAS EQUATIONS. Law and Theory. We discover facts by observation and experiment. Facts which always oc"cur under the same circumstances soon become well established. Such facts are often -summarized in a brief statement called a law. Sometimes the word law is used in the sense of the uniform behavior summarized in the brief statement. Hence, in a narrow sense, a law is a statement of a fact, but in a broad sense a law is the fact itself. Thus, the law of definite proportions (soon to be discussed) is either (i) a brief statement of the general fact of definite proportions of ele- ments in compounds, or (2) the uniform behavior itself as far as the composition of chemical compounds is concerned. The cause of many scientific facts is unknown. The explanation we give, or the statement we make, of the cause of facts is called a theory. Laws are statements of fact, theories are statements of the supposed cause of facts. Thus we know that chemical compounds have a definite composition, because we have discovered by experiment the facts on which this law is based ; and we have framed a theory, which, as far as our present knowledge is. con- cerned, is a satisfactory explanation of the cause of the general fact of definite composition. Laws seldom change, but theories are often modified. Laws are the result of experiment, theories are the outcome of mental operations. 75 76 Descriptive Chemistry. We accept a certain theory until a more satisfactory one is proposed. If a fact is not well established or is not gen- eral, we account for it by an hypothesis. An hypothesis is a guess or supposition concerning the cause of some particular fact or set of facts, and it is usually proposed as a basis for making further experiments. Hypotheses often lead to theories. Laws, theories, and hypotheses are of great service in chemistry, since they help us gather into intelligible state- ments a vast number of facts which are apparently not related. They also assist in discovering facts. Law of Definite Proportions by Weight. When the metal magnesium is heated in the air, it burns with a dazzling flame into a grayish powder, due to combination with oxygen. If a known weight of magnesium is heated in a crucible, so that the product cannot escape, a remark- able relation is revealed. In order to burn completely 1.5 gm. of magnesium, i gm. of oxygen is necessary; and the product, magnesium oxide, weighs 2.5 gm. This product contains, therefore, 60 per cent magnesium and 40 per cent oxygen. Accurate repetitions of this experi- ment have shown that this proportion by weight is fixed and definite. Again, if all the oxygen is driven from a weighed quantity of potassium chlorate by heating this compound in a crucible, 39.18 per cent of oxygen is always obtained. This means that the proportion of potassium, chlorine, and oxygen which makes up potas- sium chlorate is fixed and definite. Otherwise, the prop- erties of potassium chlorate would vary. Experiments similar to these show that in all chemical compounds the different components are always present in a definite and unvarying proportion by weight. There are no exceptions to this general fact. This constancy of proportion in Law of Multiple Proportions. 77 chemical compounds is stated as the Law of Definite Pro- portions by Weight, thus : - A given chemical compound always contains the same elements in the same proportions by weight. Sometimes it is condensed into this form : A chemical compound has a definite composition by weight. This law is one of the fundamental laws of chemistry. It is so firmly believed that if the composition of a compound is found by analysis to vary, chemists conclude that the experimental work is incorrect or that the compound is impure. The law was established as the outcome of a controversy between two French chemists, Proust (1755-1826) and Berthollet (1748-1822). The discussion lasted from 1799 to 1806. Berthollet believed that compounds might have a varying composition. Indeed, by his experiments he detected " gradual changes " in com- position. But Proust showed that Berthollet analyzed mixtures and not compounds. In a mixture the parts may be present in any propor- tion. Subsequent experiments have only strengthened our confidence in this law. Law of Multiple Proportions. Proust showed that some elements combine in more than one proportion, and thereby produce distinct compounds. But he failed to notice that if the weight of one element is constant, the varying weights of the other element are in a simple mul- tiple relation to each other. Dalton discovered this gen- eral fact about 1804. The composition of compounds is usually expressed in per cent ; but such expressions in a series of compounds reveal nothing about multiple rela- tions. If, however, a constant weight is adopted as a unit for one component, and the composition of the series of compounds is expressed in terms of this unit, then the simple multiple relation which exists between the weights of the other component is clearly seen. Thus, we learn> little from the statement that the two compounds of carbon Descriptive Chemistry. and oxygen contain 73 and 57 per cent of oxygen. But if in expressing the composition of these compounds we adopt 12 as the weight of carbon, the weights of oxygen become 32 and 16, i.e. the weights of oxygen are simple multiples. The five compounds of oxygen and nitrogen, which will soon be studied, aptly illustrate this fact : TABLE 'TO ILLUSTRATE MULTIPLE PROPORTIONS. COMPOSITION IN UNIT PER CENT. WEIGHT. RATIO. NAME. Nitrogen. Oxygen. Nitrogen. Nitroj ;en. Oxygen. Nitrous oxide .... 63.6 36,4 7 7 4 Nitric oxide j.6 6 zi A. 7 7 8 Nitrogen trioxide . . . 36.8 63.2 7 7 12 Nitrogen peroxide. . . 30.4 69.6 7 7 16 Nitrogen pentoxide . 25.9 74.1 7 7 20 From this table it is clear that the weights of oxygen combined with the same weight of nitrogen are as 1:2: 3:4:5, i.e. they are simple multiples of each other. The general fact of multiple proportions is expressed in the Law of Multiple Proportions, thus : - When two or more elements unite to form a series of compounds, a fixed weigJit of one element so combines with different weights of the other element that the relations be- tween the different weights can be expressed by small whole numbers. This law, like the law of definite proportions, is a fun- damental law of chemistry, and together they have pro- foundly influenced its theoretical and practical progress. JOHN DALTON 1766-1844 THE ENGLISH CHEMIST WHO LAID THE FOUNDATIONS OF THEORETICAL CHEMISTRY The Atomic Theory. 79 The Atomic Theory of the constitution of matter was proposed by Dalton to explain the laws of definite and multiple proportions. This theory assumes (i) that the chemical elements consist ultimately of a vast number of very small, indivisible particles or atoms, (2) that the atoms of the same element have the same weight, (3) that atoms of different elements have different weights, and (4) that chemical action is union or separation of the atoms of the elements. Let us now consider how this theory explains the facts summarized in the laws of definite and multiple propor- tions, (i) When magnesium combines with oxygen, 1.5 parts by weight of magnesium combine with one part by weight of oxygen. Analysis of the product magnesium oxide shows that this proportion is constant; that is, pure magnesium oxide always contains the elements mag- nesium and oxygen in this proportion. Now, according to the atomic theory, magnesium oxide is the product of the union of indivisible atoms of magnesium and indivisible atoms of oxygen. It therefore follows that when magne- sium and oxygen unite, atom for atom, the magnesium oxide must contain the two elements in the proportion of the weights of their atoms, i.e. it must always have the same composition. It is immaterial whether the actual weights of these elements which combine are in the pro- portion of i to 1.5, because whatever is in excess of this proportion will be left uncombined. For example, if we start with i gm. of oxygen and 2 gm. of magnesium, then 0.5 gm. of magnesium will be left uncombined. Thus the atomic theory explains the law of definite proportions. (2) But atoms do not always combine in the simple proportion of i to i. They may combine in the proportions of i to 2, 2 to 3, i to 3, i to 4, etc. But according to the atomic 8o Descriptive Chemistry. theory atoms are assumed to be indivisible. Hence, if we assume the atomic theory, the proportions of the weights of different elements in a series of compounds must be simple proportions, i.e. the elements must unite in accord- ance with the law of multiple proportions. To illustrate : There are two compounds of carbon and oxygen. Since atoms are indivisible, the simplest combinations of the atoms are (i) one atom of carbon to one atom of oxygen, and (2) one atom of carbon to two atoms of oxygen. Analysis shows that in the first compound the proportion of carbon to oxygen is 6 to 8. According to the theory, the propor- tion in the second compound should be 6 to 16; this pro- portion is verified by analysis. In other words, if we adopt 6 as the weight of carbon in its two oxides, then the weights of oxygen are in the simple proportion i to 2. Atoms and Molecules. It should not be forgotten that the laws of definite and multiple proportions deal with facts, and that the atomic theory deals with conceptions which may be true, but which cannot be proved to be true. We often speak of atoms as if they could be per- ceived by the senses, but we do so simply because such expressions help us describe, study, and interpret chemical action. According to the present views, atoms do not, as a rule, exist in the uncombined state. As soon as atoms are freed from combination, they at once unite with some other atom or atoms. The smallest particle of matter which can exist independently is not, therefore, an atom, but a group or combination of atoms. These groups of atoms are called molecules. If the atoms in a molecule are atoms of the same element, then the molecule is a molecule of an element; but if the atoms of different elements are combined, then the molecule is the molecule of a compound. All matter, as a rule, consists of mole- Chemical Symbols. 8 cules, and the molecules are made up of atoms. A mole- cule of a few elements contains only one atom. Chemists define a molecule as the smallest part of a compound or of an element which can exist in the free state and mani- fest the properties of the compound. Thus, the smallest particle of water is a molecule of water, but a molecule of water contains smaller particles still, viz., atoms of hydro- gen and oxygen. We may define an atom as the indivis- ible constituent of a molecule. It is also the smallest particle of an element which takes part in chemical changes. Our views regarding molecules are based on extensive study of the physical properties of gases. The molecule is often spoken of as the physical unit, because in physical changes molecules are not decomposed. Whereas the atom is the chemical unit, because it enters into all chemi- cal action. The molecule is chemically divisible, but the atom is chemically indivisible. Chemical Symbols, which were mentioned in Chapter I, are designed to represent single atoms. Thus, H repre- sents one atom of hydrogen, O one atom of oxygen, N one atom of nitrogen. If more than one atom is to be desig- nated, the proper numeral is placed before the symbol, 2 H means 2 atoms of hydrogen. 3 O means 3 atoms of oxygen. 4 P means 4 atoms of phosphorus. But if the atoms are in chemical combination, either with themselves or with other atoms, then a small numeral is placed after and a little below the symbol, thus : H 2 means 2 atoms of hydrogen in combination, N 3 means 3 atoms of nitrogen in combination, P 4 means 4 atoms of phosphorus in combination. 8 2 Descriptive Chemistry. Chemical Formulas. A formula is a group of symbols which is designed to express the composition of a com- pound. In writing a formula the symbols of the different atoms making up the compound are placed side by side. Thus, H 2 O is the formula of water, because this group of symbols is the simplest expression of the facts which are known about this compound. Similarly, KC1O 3 is the formula of potassium chlorate. These symbols might be written in a different order, but usage has determined the order in this, as in most cases. A formula represents one molecule. Hence, KC1O 3 represents one molecule of potassium chlorate, and means that the molecule of this compound contains one atom each of potassium and chlo- rine and three atoms of oxygen. If we wish to designate several molecules, the proper numeral is placed before the formula, thus : 2 KC1O 3 means 2 molecules of potassium chlorate. 3 H 2 O means 3 molecules of water. 4 H 2 SO 4 means 4 molecules of sulphuric acid. In certain compounds some of the atoms act like a single atom in chemical changes. This fact is often expressed by inclosing the group of atoms in a parenthesis, or by sepa- rating it from the rest of the formula by a period. Thus, the formula of ammonium nitrate is (NH 4 )NO 3 . Simi- larly, the formula of alcohol is often written C 2 H 5 . OH, because the groups C 2 H 5 and OH act as units. The use of the period is confined mainly to organic and mineralogi- cal chemistry. It is sometimes omitted, especially if the composition of the compound is well understood. If a group of atoms is to be multiplied, it is placed within a parenthesis. Thus, the formula of lead nitrate is Pb(NO 3 ) 2 . This means that the group NO 3 is to be multiplied by 2. Chemical Equations. 83 The formula 2 Pb(NO 3 ) 2 means that the whole formula is to be multiplied by 2. Symbols and formulas are sometimes used to represent an indefinite amount of an element or compound. Thus, O may mean oxygen and H.jSO 4 sulphuric acid, regardless of the amount. This use of symbols and formulas saves time, but it is not scientific. They are often thus used to label bottles in a laboratory. Such a departure from accuracy should not be allowed to obscure their real meaning. The complete significance ot symbols and formulas can be grasped only by their intelligent use. They should not be committed to mem- ory slavishly. It is desirable, however, to learn the common ones while the substances they represent are being studied, and consider their relations more fully when the needed facts have accumulated. (See Chapters IX and XIII.) A Chemical Reaction is a special or limited chemical change. When potassium chlorate is heated, the chemical change results finally in the liberation of all the oxygen and the formation of potassium chloride. Such a change is called the reaction for preparing oxygen from potassium chlorate, or the reaction for the decomposition of potas- sium chlorate. Obviously, the study of chemistry is largely a study of reactions. Chemical Equations. In expressing various facts about chemical reactions, it is customary to use an equa- tion consisting of the proper symbols or formulas. Sub- stances entering into the initial stage of a reaction are called factors, and those present in the final stage are called products. The symbols and formulas of the factors connected by the sign plus ( -f- ) are placed at the left of the sign of equality, and those of the products at the right. Equations are usually read from left to right. Occasion- ally the words reaction and equation are used as synonyms, but such a use is inaccurate and confusing. 84 Descriptive Chemistry. When magnesium burns in the air or in oxygen, mag- nesium oxide is formed. The simplest equation for this reaction is Mg + O = MgO Magnesium ' Oxygen Magnesium Oxide This equation is read : Magnesium and oxygen form mag- nesium oxide. It means, also, that when magnesium and oxygen' react, one atom of magnesium unites with one atom of oxygen and forms one molecule of magnesium oxide. The simplest equation for the preparation of hy- drogen by the reaction of zinc and sulphuric acid is Zn+ H 2 SO 4 = H 2 + ZnSO 4 Zinc Sulphuric Acid Hydrogen Zinc Sulphate This equation is read: Zinc and sulphuric acid form (or produce) hydrogen and zinc sulphate. It means, further, that one atom of zinc and one molecule of sulphuric acid form one molecule (or two atoms) of hydrogen and one molecule of zinc sulphate. By similar equations we may express certain facts about all reactions which are under- stood. The above equations might be called ordinary chemical equations, or atomic equations. Other forms are used, and they will be discussed in Chapters IX, X, and XIII. The following facts about ordinary chemical equations should be noted : (1) The sign plus does not necessarily mean addition chemically. It does in the equation Mg -f O = MgO, but not in the equation HgO = Hg-fO. In the latter the products are merely mixed. The sign plus may be expressed by the words and* acted upon, added to, mixed with. The sign equality is often read equal, give, form, or produce. (2) Equations do not always include all the participating substances. In Mg + O = MgO no nitrogen (N) appears because nitrogen takes no Exercises. 85 chemical part in the change, despite the fact that the air is largely nitrogen. Similarly, in Zn + H 2 SO 4 = H 2 + ZnSO 4 , no water (H 2 O) appears, because the water (in the dilute sulphuric acid) simply serves to dissolve the zinc sulphate from the surface of the zinc. A special form of equation, called the ionic equation, is used to express chemical changes which occur in solution (see Chapter X). (3) Equations tell nothing about the heat changes (see Chapter X). (4) Most equations represent only the beginning and end of reac- tions. Thus, in KC1O 3 = O 3 + KC1 several changes do not appear, because the purpose of this equation is to express the complete decom- position of potassium chlorate nothing else. EXERCISES. 1. Define law, theory, and hypothesis as used in science. 2. State the law of definite proportions. Illustrate it. Give a brief account of its discovery. 3. State the law of multiple proportions. Illustrate it. Who dis- covered it? When? 4. State the atomic theory. What are atoms according to this theory? How are atoms related to chemical action? How are atoms related to molecules? What is a molecule? 5. What is the symbol of an element? How are they formed? Interpret the symbols : H, 2O, N 3 , 2 P, 30, K 2 , S 2 , 2 Cl. 6. What is the formula of a compound? What does a formula represent? Interpret the formulas: H 2 O, 2 H 2 O, KC1O 3 , 4 H 2 SO 4 , (NH 4 )NO 3 , C 2 H 5 .OH, Pb(N(X) 2 , Ca(OH) 2 . " 7. Give the symbols of the following elements : oxygen, hydrogen, nitrogen, zinc, copper, magnesium, platinum, iron, sodium, sulphur, carbon, mercury. 8. What elements correspond to the following symbols : Na, Cu, K, Zn, S, P, Pt, Pb, H, Hg, Fe, Mg? 9. Give the formulas of the following compounds : water, potas- sium chlorate, sulphuric acid, magnesium oxide. 10. Define and illustrate the term chemical reaction. 11. What is a chemical equation ? For what is it used? What are factors and products in an equation? How are equations written? Illustrate your answer. How are they read ? 86 Descriptive Chemistry. 12. Interpret the equation : Mg + O = MgO. 13. What does the plus ( + ) sign mean in the above equation? What other meanings has this sign? 14. State several facts about equations. PROBLEMS. 1. How many centigrams in 1745 kg.? In 250 gm.? In 1425 dg. ? 2. How many cubic centimeters in 50 1. ? In I cu. dm. ? 3. What is the weight of (a) loocc. of hydrogen, and (6) 25 1. of oxygen, under standard conditions ? 4. What weight of (a) hydrogen and (<) oxygen can be obtained from 1 80 gm. of water ? 5. What (#) weight and ($) volume of oxygen are necessary to unite with 200 kg. of hydrogen ? 6. What weight of hydrogen is necessary to unite with the oxygen in 100 gm. of air to form water ? (Assume that air is one fifth oxygen.) CHAPTER VIII. ACIDS, BASES, AND SALTS. Introduction. Many chemical compounds fall naturally into one of three groups, long known as acids, bases, and salts. Not all compounds, of course, are included in this classification. Each group has its characteristic properties, 'though the groups are closely related and sometimes over- lap. Many familiar substances belong to these groups. A knowledge of the properties of acids, bases, and salts,, of their special behavior, and of their intimate relations is essential in the study of chemistry. General Properties of Acids, Bases, and Salts. Acids have a sour taste. The early chemists detected this property, and the word acid (from the Latin acidtts, sour) emphasizes the fact. Acids change the color of many vegetable substances. Thus, blue litmus is turned red by acids. Acids also have the power to decompose most carbonates, like limestone, thereby liberating carbon diox- ide gas which escapes with effervescence. Most bases have a slimy, soapy feeling, and a bitter taste. They turn red litmus blue. Caustic soda and ammonium hydroxide are bases. Many salts have the well-known salty taste. Sodium chloride, the familiar table salt, is an example. Usually, they have no action on litmus. All acids contain hydrogen, which is usually liberated when metals and acids interact. Most acids contain oxy- gen. For many years it was thought that oxygen was an 87 88 Descriptive Chemistry. essential component of all acids, and its name, oxygen (derived from Greek words meaning " acid producer ") was given by Lavoisier because of this belief (see Discovery of Oxygen). We now know that hydrogen, not oxygen, is the essential component of all acids. Another necessary component of acids is some element like nitrogen, sulphur, chlorine, or phosphorus, which belongs to a class of elements called non-metals. For this reason it is some- times convenient to think of non-metals as the elements which form acids. Thus sulphuric acid contains sulphur, besides hydrogen and oxygen ; while hydrochloric acid contains chlorine, besides hydrogen. Bases contain oxygen and usually hydrogen, but their distinctive component is a metal, e.g. sodium, potassium, calcium. Hence a metal may be properly regarded not merely as an element possessing in a varying degree the physical properties of hardness, luster, power to conduct heat and electricity, but also the chemical property of forming bases. Salts contain a metal and a non-metal, and most of them contain oxygen. Thus, potassium nitrate contains the metal potassium and the non-metal nitrogen, besides oxygen ; while potassium chloride contains potassium and the non-metal chlorine, but no oxygen. The nature o*f acids, bases, and salts is clearly shown by their chemical relations to each other. When acids and bases interact, salts are formed. That is, the acid and base destroy more or less completely the marked prop- erties of each other and produce a compound which has few, and often none, of the properties of the original acid or base. The acid and base neutralize each other. An example will make this point clear. When hydrochloric Acids, Bases, and Salts. 89 acid and sodium hydroxide interact, sodium chloride and water are formed. The chemical change may be written thus- HC1 + NaOH. NaCl + H 2 O Hydrochloric Acid Sodium Hydroxide Sodium Chloride Water This equation represents the facts which have been repeatedly verified by experiment. This series of chemi- cal changes is called neutralization, and later it will be more fully discussed. Taking this equation as a type of the chemical changes which occur in neutralization, it is clear that in such changes, generally speaking (i) the metal of the base takes the place of the hydrogen of the acid, thereby forming a salt, while (2) the hydrogen of the acid combines with the hydrogen and oxygen of the base to form water. In neutralization the hydrogen and oxygen of the base act as a unit. This group of atoms (OH) is called hydroxyl. Compounds containing this group are called hydroxides. Hydroxyl does not exist free and uncombined like elements and compounds, but it acts like a single atom in many changes. It is called a radical. To emphasize the fact that it is a unit, the hydroxyl group is sometimes put in a parenthesis, e.g. Ca(OH) 2 . Hydroxides are often said to be founded on the water type. Thus we have Water HOH Sodium hydroxide . . . . ' NaOH Potassium hydroxide .... KOH Calcium hydroxide .... Ca(OH) 2 Hence we may regard sodium hydroxide and potassium hydroxide as water in which the hydrogen atom has been replaced by a metallic atom. The words hydroxide, hydrate, and hydroxyl are all derived from hudor, the Greek word for water. 90 Descriptive Chemistry. The most characteristic property of acids and bases is, then, this power to neutralize each other and thereby form salts and water. Acids. The common acids are sulphuric acid, hydro- chloric acid, nitric acid, and acetic acid. Many acids are liquid, as sulphuric and nitric ; a few are gases, as hydro- chloric ; others are solid, as tartaric, citric, oxalic. Most are soluble in water, and such solutions are familiarly called acids. These solutions may be dilute or concen- trated, and the general properties vary somewhat with the strength. Concentrated acids are usually corrosive and should be handled with precaution, even when one is thoroughly familiar with their properties. Substances which turn blue litmus to red are said to contain an acid, to be acid, or to have an acid reaction. The exact nature, however, of such a substance must be determined by additional tests. Many familiar substances are acids or contain them. Vinegar, pickles, and similar relishes contain dilute acetic acid. Lemon juice is mainly citric acid. Sour milk con- tains lactic acid. Unripe fruits, sour bread, and sour wines contain acids. " Soda water " is a solution of carbonic acid (or more accurately carbon dioxide), and " acid phosphate" is a solution of a sour calcium phosphate. No brief, satisfactory definition of an acid can be given, for chemists do not agree on this point. We might say, however, that an acid is a compound containing hydrogen which can be replaced by a metal; but this definition includes water, since its hydrogen is readily replaced by sodium. Not only must the hydrogen of an acid be replaced by a metal, but one product of the reaction must be a salt. The replacing metal may, of course, come from a compound, e.g. an oxide, hydroxide, or carbonate. Acids, Bases, and Salts. 91 Nomenclature of Acids. Oxygen is a component of most acids, and the names of these acids correspond to the proportion of oxygen which they contain. The best known acid of an element usually has the suffix -ic, e.g. sulphuric, nitric, phosphoric. If an element forms another acid, containing less oxygen, this acid has the suffix -ous, e.g. sulphurous, chlorous, phosphorous. Some elements form an acid containing less oxygen than the -ous acid ; these acids retain the suffix -ous, and have, also, the prefix hypo-, e.g. hyposulphurous, hypophosphorous, hypochlo- rodl. Hypo- means under or lesser. If an element forms an acid containing more oxygen than the -ic acid, such an acid retains the suffix -ic, and has, also, the prefix per-, e.g. persulphuric, perchloric. The prefix per- means beyond or over. The few acids which contain no oxygen have the prefix hydro- and the suffix -ic, e.g. hydrochloric, hydrobron\i, hydrofluoric. It should be noticed that these suffixe^ are not always added to the name of the element, but often to some modification of it. The nomenclature of acids is well illustrated by the series of chlorine acids : * ACIDS OF THE ELEMENT CHLORINE. NAME. FORMULA. Hydrochloric Hypochlorous Chlorous HC1 HC1O HC1O 2 Chloric Perchloric HC10 3 HC1O 4 Not all elements form a complete series of acids, but the nomenclature usually agrees with the above principles. 92 Descriptive Chemistry. Some acids have commercial names. Thus, sulphuric acid is often called oil of vitriol, and hydrochloric acid is known as muriatic acid. Acids in which carbon is the essential component end hi -ic, but they are often arbitrarily named (see Organic Acids). An examination of the formulas of acids shows that all do not con- tain the same number of hydrogen atoms. Acids are sometimes classi- fied by the number of hydrogen atoms which can be replaced by a metal. This varying power of replaceability is called basicity. A monobasic acid contains only one atom of replaceable hydrogen in a molecule, e.g. nitric acid, HNO 3 . A molecule of acetic acid (C 2 H 4 O 2 ) contains four atoms of hydrogen, but for reasons which are too complex to state here, only one of these atoms can be replaced by a metal. Dibasic and tribasic acids contain two and three replaceable hydrogen atoms, e.g. sulphuric acid (H 2 SO 4 ) and phosphoric acid (H 3 PO 4 ). Obviously, monobasic acids form only one class of salts, dibasic acids form two classes, tribasic acids form three, and so on. Bases. The term base, in a narrow sense, means the strong bases, which are very soluble in water, and are com- monly known as alkalies, e.g. sodium, potassium, and ammonium hydroxides. In a broad sense it means any substance which will neutralize an acid, e.jr. calcium oxide, ammonia gas, as well as the hydroxides of metals. Most bases are solids ; but since they are usually soluble in water, these solutions, as in the case of acids, are familiarly called the base, or alkali, itself. Concentrated alkalies, like concentrated acids, are corrosive. The common alka- lies sodium and potassium hydroxides are often called caustic soda and caustic potash to emphasize this property ; and calcium oxide, or lime, is sometimes called caustic lime ; the corrosive nature of ammonium hydroxide, or ordinary 'ammonia, is also well known. Substances which turn red litmus to blue are said to contain an alkali (or base), to be alkaline, or to have an alkaline reaction. Acids, Bases, and Salts. 93 The word basic is often used instead of alkaline. Other tests besides that with litmus must be applied, however, to determine the exact nature of a substance having an alka- line reaction. Alkalies dissolve grease and fats, and are often used as cleansing agents, ammonium hydroxide being widely employed for this purpose. They also inter- act with fats to form soaps, large quantities of sodium hydroxide being annually utilized in the soap industry (see Soap). A base, like an acid, is rather difficult to define. We might say that a base is an hydroxide or oxide of a metal, which will neutralize an acid, thereby forming a salt. The term must include ammonia, which does not contain a metal. But, as we shall see later, a certain combination of elements related to ammonia acts like a metal (see Ammonium). Nomenclature of Bases. There is no general rule covering the nomenclature of bases, as in the case of acids. Since most bases contain hydrogen and oxygen, they are often called hydroxides. Hydrate is sometimes used as a synonym of hydroxide. The term alkali em- phasizes general properties rather than suggests specific composition. Hydroxides are distinguished from each other by placing the name of the metal before the word hydroxide, e.g. sodium hydroxide, potassium hydroxide, calcium hydroxide. The common hydroxides have long been known by several names. Thus, calcium hydroxide is often called limewater. Ammonium hydroxide is some- times called ammonia water or simply (but inaccurately) ammonia, and it was formerly called volatile alkali. Be- sides the common names of the hydroxides of sodium and potassium already given, they are sometimes called fixed alkalies. 94 Descriptive Chemistry. Not all bases contain the same number of hydroxyl groups. Hence bases, like acids, may form one or more salts. This power is called acidity. Bases are called monacid, diacid, triacid bases, etc., accord- ing to the number of replaceable hydroxyl groups present in a molecule. Thus, calcium hydroxide (Ca(OH) 2 ) is a diacid base, and aluminium hydroxide (A1(OH) 3 ) is a triacid base. Salts. Sodium chloride, or ordinary table salt, is the most familiar salt. It has been known for ages. Doubt- less this class of chemical compounds received its name because of the general resemblance most of them bear to common salt. Most salts are solid and are soluble in water. Many of them have no action on litmus, and are, therefore, said to be neutral or to have a neutral reaction. This indifference to litmus is not a decisive test for a salt, since many other substances, water for example, have no action on litmus. Nevertheless the term neutral is applied to substances which do not change the color of litmus. Some substances which are salts, as far as their structure and method of formation are concerned, do not have a neutral reaction. Thus, sodium carbonate, which is the sodium salt of carbonic acid, has a marked alkaline reaction, being in fact known in commerce simply as " alkali." A salt may be defined as the main product of the inter- action of an acid and a base. It may, however, be a sub- stance which has the properties of a salt, regardless of the method of formation. Salts are formed in various ways. The interaction of an acid and a base has been mentioned. The interaction of acids with oxides of cer- tain metals or with metals themselves produces salts. Sodium oxide and sulphuric acid interact and form the salt sodium sulphate, thus : Na,O + H 2 SO 4 Na 2 S0 4 + H 2 O Sodium Oxide Sulphuric Acid Sodium Sulphate Water Acids, Bases, and Salts. 95 While zinc and sulphuric acid, as already stated, form the salt zinc sulphate as well as hydrogen, thus : Zn + H 2 S0 4 = ZnS0 4 + H 2 Zinc Sulphuric Acid Zinc Sulphate Hydrogen Carbonates interact with acids and form other salts. Calcium carbonate and hydrochloric acid form the salt calcium chloride, thus : CaC0 3 + 2HC1 = CaCl 2 + CO 2 + H 2 O Calcium Hydrochloric Calcium Carbon Water Carbonate Acid Chloride Dioxide Nomenclature of Salts. The name of salts containing oxygen are derived from the name of the corresponding acid. The characteristic suffix of the acid is changed to indicate this relation. Thus, the suffix -ic becomes -ate, and the suffix -ous, becomes -ite. Hence : Sulphuric acid forms sulphates. Sulphurous acid forms sulphites. Nitric acid forms nitrates. Nitrous acid forms nitrites. Chloric acid forms chlorates. Hypochlorous acid forms hypochlorites. Permanganic acid forms permanganates. $ The name of the replacing metal is retained, e.g. potas- sium chlorate, sodium sulphate, calcium hypochlorite, po- tassium permanganate. Notice that the prefixes hypo- and per- are not changed. The names of salts containing only two elements, fol- lowing the general rule for binary compounds, end in -ide. This suffix is added to a modification of the name of the non-metal, giving the names chloride, bromide, sulphide, fluoride, etc. The prefix hydro- which is contained in the 96 Descriptive Chemistry. name of the acid is omitted. Thus, the name of the sodium salt of hydrochloric acid is sodium chloride ; simi- larly, there are the names potassium chloride, calcium fluoride, and sodium iodide. Sometimes, the salts of these hydrogen acids are called halides to emphasize their rela- tion to common salt, which in Greek is called halos. Salts in which all the hydrogen atoms of the corresponding acid have been replaced by a metal are called normal salts, e.g. sodium sulphate, Na.,SO 4 . If some of the hydrogen atoms are not replaced by a metal, an acid salt is formed. Thus, acid sodium sulphate may be regarded as derived from sulphuric acid, which is dibasic, by replacing one of the atoms of hydrogen by sodium, though of course the salt is not prepared in this way. Expressed as formulas these relations may be written thus : Acid Acid Salt Normal Salt H,SO 4 HNaSO 4 Na 2 SO 4 Only those acids which contain two or more replaceable hydrogen atoms form acid salts. On the other hand, if not all the hydroxyl groups of a base are replaced when the base reacts with an acid, then a basic salt results. Thus, basic nitrate of bismuth may be regarded as the salt derived from bismuth hydroxide (Bi(OH) 3 ) by replacing one hydroxyl group of the base by the group NO 3 of nitric acid. The formula of this basic nitrate of bismuth is Bi(OH) 2 NO 3 . The following equation illustrates the changes : Bi(OH) 3 + HN0 3 = Bi(OH) 2 NOo + H 2 O Bismuth Hydroxide Nitric Acid Basic Bismuth Nitrate Water Only those bases having two or more hydroxyl groups can form basic salts. Some basic salts are very complex. Relation of Oxides to Acids and Bases. Most non- metallic elements form oxides which unite with water and produce an acid. The oxides of many metallic elements, Acids, Bases, and Salts. 97 on the other hand, unite with water and produce hydrox- ides. The two oxides of the non-metal sulphur act thus 50 2 + H 2 O = H 2 SO 3 Sulphur dioxide Water Sulphurous Acid 50 3 + H 2 O = H 2 SO 4 Sulphur Trioxide Water Sulphuric Acid The oxide of the metal calcium acts thus CaO -f H 2 O = Ca(OH) 2 Calcium Oxide Water Calcium Hydroxide Oxides of non-metals which unite with water and thereby produce acids are called anhydrides, i.e. literally, sub- stances without water. Examples are carbonic anhydride (CO 2 ), sulphuric anhydride (SO 3 ), phosphoric anhydride (P 2 O 5 ). Oxides of metals which produce hydroxides are called basic oxides. A few oxides behave exceptionally. It is convenient to regard an anhydride as the root or basis of its corresponding acid, and a basic oxide as the root of its hydroxide. The fact that many non-metallic oxides redden moist blue litmus led Lavoisier into the erroneous belief that oxygen is an essential compo- nent of acids. And some authorities even now (incorrectly) speak of these oxides as acids ; thus, carbon dioxide (CO 2 ) is occasionally called carbonic acid. The compounds which Lavoisier galled acids were anhy- drides. And it was not until about 181 1 that Davy showed (i) that some acids do not contain oxygen (e.g. hydrochloric acid, HC1 ), and (2) that the so-called acids of Lavoisier are not real acids until they have obtained hydrogen from the water with which they combine. Neutralization has been defined as the series of changes whereby acids and bases mutually destroy each other's characteristic properties and produce a salt and water. 98 Descriptive Chemistry. But neutralization has a deeper meaning and broader ap- plication than the mere destruction of properties. If measured volumes of different acids are exactly neutralized by different alkalies, remarkable relations are revealed. This may be done by dropping one into the other from a graduated tube, called a burette (Fig. 15). The exact point of neutralization is shown by an indicator; this is a solution of litmus or some other substance, which tells by the color whether the solution is acid or alkaline. Experiment shows that (i) a definite quan- tity of an acid neutralizes a definite quantity of an alkali, (2) the same acid is neutralized by different quantities of different alkalies, and (3) the ratio of the quantities of the FIG. 15. Burettes. different alkalies is the same for all acids. 1 EXERCISES. 1. Define and illustrate (a) an acid, (6) a base, (c) a salt, (d} an al- kali, (e) hydroxyl, (/) an hydroxide. 2. Name three common acids and bases. State the general proper- ties of each class. 3. Define and illustrate (a) neutralization, (b} acidity of bases, (c} basicity of acids, ( K = 39.06. 6. How much oxygen can be prepared from (ft) 122.5 m< f potas- sium chlorate, (b) 245 gm., and (c) 421 gm. ? Solution. The equation is KC10 3 = 3 + KC1 122.5 =48 4- 74-5 These equation weights are obtained by adding the atomic weights found in the table, (a) By inspection, 122.5 g m - of potassium chlorate yield 48 gm. of oxygen. (b) The proportion needed is 122.5 : 48 :: 245 : x. And x = 96 gm. (c) Similarly, 122.5 : 4^ : : 4 21 ' * And x 164.9. 7. (a) How much oxygen can be prepared from 50 gm. of potassium chlorate, and (b) how much potassium chloride will remain? Ans. (a} = 19.59, (b) = 30.41. 8. A certain weight of potassium chlorate was heated until completely decomposed. The residue weighed 20.246 gm. (a) What was its weight ? (b) How much oxygen was evolved? Ans. (a)= 33.29, (b) 13.044. 9. What weight of potassium chlorate is needed to generate 144 gm. of oxygen? .^^.367.5. 10. What weight of potassium chloride remains after obtaining 8 gm. of oxygen from potassium chlorate? Am. 12.416. 1 1 . How many grams of oxygen can be generated from 490 gm. of po- tassium chlorate? Ans. 192. 12. How much hydrogen can be prepared from (a) 65 gm. of zinc, (b) 130 gm., (V) 297 gm.? Ans. (c} 9.14- 13. How much zinc is needed to prepare (a) 2 gm. of hydrogen, (b) 14 gm., and (c) 17 gm.? 14. How much zinc sulphate can be prepared from (a) 98 gm. of sul- phuric acid, (b) 196 gm., and (c) 427 gm.? Ans. (c) 701.5. 15. A balloon holds 132.74 kg. of hydrogen. How much (a) zinc and (b) sulphuric acid are needed to produce the gas? Ans. (a) 43 I 4-05> (*) 6504.26. no Descriptive Chemistry. 1 6. How much (a) mercury and () oxygen can be obtained from 10 gm. of mercuric oxide? (Equation is HgO = Hg + O, or 216 = 200 + 1 6.) Ans. (a) 9.259, (b} 0.74. 17. How much mercury will remain after obtaining 48 gm. of oxygen by heating mercuric oxide? 1 8. A lump of carbon weighing 24 gm. is burned in air. What weight of (a) carbon dioxide is formed and () oxygen is needed? (c) If a liter of oxygen weighs 1 .43 gm., what volume of oxygen is needed ? (Equation is C + O 2 = CO 2 , or 12 + 32 = 44.) Ans. (c) 44.75 1. 19. What weight of carbon dioxide is formed by burning 112 Ib. of coal containing 15 per cent of impurities? 20. A lump of sulphur weighing 32 gm. is burned in air. Calculate the weight of (a) oxygen needed and (&) sulphur dioxide formed. (Equation is S + O 2 = SO 2 , or 32 + 32 = 64.) 21. Calculate the weight of oxygen needed to burn 731 gin. of sul- phur containing 15 per cent of impurities. Ans. 621.35. 22. What weight of sulphur dioxide is formed by burning 67 per cent of 8794 kg. of sulphur? CHAPTER X. LIGHT, HEAT, ELECTRICITY, AND CHEMICAL ACTION. CHEMICAL action is always manifested by one or more of the different forms of energy, such as light, heat, and electricity. This means that a chemical change involves not only a rearrangement of matter, but also a transfor- mation of energy. Thus, when coal is burned, a new compound called carbon dioxide is formed, but heat is also liberated. Sometimes we pay more attention to the result- ing matter than to the energy, but both are involved. In the present chapter we shall emphasize the relation of energy to chemical action. The law of the conservation of energy should be recalled in this connection. Energy, like matter, cannot be created or destroyed ; we can only transform it. And the transformation involves no loss or gain. Hence, chemical energy, which is the immediate cause of chemical action, will appear as heat, light, or electricity. The Relation of Light to Chemical Action is illustrated in photogra- phy. Coatings consisting of compounds of silver and organic matter are quickly blackened by light (see Photography). Sunlight fades many colors. It likewise assists the chemical changes involved in the growth of plants. The formation of the green coloring matter of foliage is partly due to sunlight. A mixture of hydrogen and chlorine gases re- mains unchanged in the dark, but in direct sunlight it explodes violently. On the other hand, light is often a product of chemical action. Many chemical experiments show this, especially those with oxygen. Sparks, most flames, and the flash of a gun are other illustrations of the close relation between light and chemical action. Combustion in its varied forms is also manifested by light, as well as by heat. HI H2 Descriptive Chemistry. HEAT AND CHEMICAL ACTION. Heat and Chemical Action are closely and definitely related. Every chemical change is attended by the libera- tion or absorption of heat. Moreover, the heat involved can often be measured. Heat is measured in calories, a calorie being the quantity of heat necessary to raise the temperature of I gm. of water from o to i C. For example, the heat liberated by the burning of I gm. of hydrogen is 34,200 cal., and of I gm. of pure charcoal is about 8000 cal. Attention has already been called to the high temperature of the hydrogen flame (see Chapter III). Ordinary chemical equations do not express changes in energy. To represent heat changes, the number of calories of heat involved is placed after the equation, thus : H 2 + O = H 2 O + 68,400 cal. Hydrogen Oxygen Water This is called a thermal equation, and it means that 68,400 cal. of heat are liberated, when 2 gm. of hydrogen unite with 16 gm. of oxygen to form 18 gm. of water. In some changes heat disappears. Thus, when carbon unites with sulphur to form carbon disulphide, heat is absorbed. The equation expressing this fact is C + S 2 = CS 2 19,600 cal. Carbon Sulphur Carbon Disulphide Heat involved in the formation of a particular compound is called heat of formation of that compound. If heat is liberated in the formation of a compound, the heat is called positive ( + ); and the compound is termed exothermic. Heat of formation which is absorbed is called negative ( ) ; and a compound having a negative heat of formation is said to be endothermic. Exothermic compounds are stable, and can be decomposed only by the addition of the same quantity of heat liber- ated by their formation. Thus, 68,400 cal. of heat, or an equivalent quantity of energy, must be added to 18 gm. of water to decompose it Light, Heat, Electricity, and Chemical Action. 113 into 2 gm. of hydrogen and 16 gm. of oxygen. Such heat is called heat of decomposition. On the other hand, endothermic compounds are unstable, and often explosive. They decompose easily with the liberation of heat. Ozone is endothermic. Heat is absorbed during its formation from oxygen ; but when ozone decomposes, heat is liber- ated. Two parts (by volume) of ozone form three parts (by volume) of oxygen and liberate 72,400 cal. A familiar instance of the evolution of heat by chemical action is the slaking of lime. When lime and water are mixed, their union produces sufficient heat to boil water and often to set fire to wood. Steam can be seen escaping from the boxes in which lime is being mixed with water and sand to form plaster or mortar. Buildings in which lime is stored sometimes take fire, if rain leaks in upon the lime. Ships loaded with lime are in constant danger of being burned. Other substances liberate heat when added to water, e.g. sulphuric acid, sodium and potassium hydrox- ides, and the metals, sodium and potassium. Heat is the initial cause of many chemical changes. It is necessary to start many reactions, just as a stone on top of a hill must be pushed before it will roll toward the bottom. Hydrogen and oxygen mix freely without com- bining, but union occurs the instant heat is applied in form of a flame or an electric spark. Similarly, illuminating gas must be lighted, i.e. raised to the kindling tempera- ture before the chemical changes which cause the light and heat can proceed. These facts mean that chemical action often depends upon temperature. This statement has been strikingly illustrated in the last four years. At the extremely low temperature obtained by using liquid air and similar substances, it appears that many chemical reactions cease. While at the exceedingly high temperature pro- duced by electricity many changes, chemical and physical, hitherto impossible, occur quickly and simply. 114 Descriptive Chemistry. The Electric Furnace of Moissan. Until recently the heat needed for chemical changes was obtained by burn- ing carbon or its compounds, such as charcoal, illumi- nating gas, and oil. Sometimes the blast lamp and oxyhydrogen blowpipe were used. But all these sources have been surpassed in efficiency by the electric furnace. It is well known that an electric arc light produces in- tense heat. The high temperature of the arc, i.e. space between the glowing ends of the carbons, is unequaled by that of any other source of artificial heat. If the carbon rods are inclosed in a box that prevents the escape of heat, a temperature estimated to be about 3500 C. is produced inside the box. This apparatus is called an electric furnace. It was devised and perfected by the French chemist, Mois- san, and used by him in experimenting at high temperatures. One form of the electric furnace is shown in Figure 16. FIG. 16. Moissan's electric furnace. Moissan's description of this furnace is as follows : " It consisted of two bricks of quicklime placed one on top of the other. The lower brick contained a longitudinal groove to receive the two electrodes [carbon rods], and situated in the center was a small cavity. This cavity might vary in size, and contained a bed some centimeters in depth of the substance to be acted upon by the heat of the arc, or a small crucible of carbon containing the substance to be treated may be placed there. The upper brick was slightly hollowed out in the part just above the arc. As the intense heat of the current soon melted the HENRI MOISSAN 1852 THE EMINENT FRENCH CHEMIST WHOSE DISCOVERIES CONTINUE TO ENRICH INORGANIC CHEMISTRY Light, Heat, Electricity, and Chemical Action. 115 surface of the lime, giving it, at the same time, a beautiful polish, a dome was obtained in this way which reflected all the heat on to the small cavity which contained the crucible." Figure 17 is a vertical sec- tion of the furnace, showing the parts slightly separated. The furnace is small, some being only 16 to 18 cm. (about 7 in.) long, 15 cm. wide, and 14 cm. high. The carbon rods are from i to 5 cm. in diameter. When a cur- rent is passed through the Car- FIG. 17. Vertical section of Moissan's electric furnace. bon rods, the tremendous heat produced is retained in the space by the non-conducting walls and acts upon the substance below the arc. The outside of the furnace remains cold enough to be touched by the hand, but the inside is almost twice as hot as the oxyhydrogen flame. There is no electrical action upon the chemicals. The intense heat alone pro- duces the remarkable changes, which are often accom- plished in a few minutes. Sand, lime, magnesium oxide, and other refractory oxides melt and volatilize. The ele- ments carbon, silicon, and boron boil; and gold, copper, and platinum quickly melt and vaporize. Large masses of rare and uncommon elements are quickly reduced from their oxides and obtained in the pure state, e.g. chromium, manganese, tungsten, uranium, and molybdenum. Char- coal becomes graphite. And stable compounds of carbon, boron, and silicon are formed. These are the carbides, borides, and silicides. Some of the carbides have an in- dustrial use as well as scientific interest, especially calcium carbide and silicon carbide (see below). Other carbides are the sources of pure metals, since the fusion of a car- bide and oxide of the same metal yields the metal itself. ii6 Descriptive Chemistry. Industrial Use of the Electric Furnace. Huge elec- tric furnaces constructed on the type devised by Moissan are in active operation. And since electricity is now ob- tained in many localities by running dynamos by water, new industries requiring intense and continuous heat have recently sprung into existence. Several of these plants are located at Niagara Falls, which furnishes enormous power at a relatively small expense. Calcium Carbide is made on a large scale by heating a mixture of lime and coke (a form of carbon) in an electric furnace. The chemical change is caused solely by the intense heat and may be represented thus : 3C + CaO = CaC 2 + CO Carbon Lime Calcium Carbide Carbon Monoxide This method of making calcium carbide cheaply was dis- covered independently and at about the same time (1892- 1895) by Moissan and Willson. The furnaces now in operation vary in details, but all have one essential feature, viz., the heat is generated by an electric current passing between two carbon electrodes. In most furnaces one elec- trode is a crucible wholly or partly of carbon, and the other electrode is a stout carbon pillar dipping into the mixture. Calcium carbide is a hard, brittle, dark gray, crystalline solid with a metallic luster. Its specific gravity is 2.2. The most striking and useful property is its action with water, acetylene being formed, thus: CaC 2 + 2H 2 = C 2 H 2 -f Ca(OH) 2 Calcium Carbide Water Acetylene Calcium Hydroxide Calcium carbide is used to generate acetylene gas. This gas burns with a brilliant flame, and is coming into general use as an illuminant. Owing to its action with water, Light, Heat, Electricity, and Chemical Action. 117 calcium carbide is packed and sold in air-tight cans (see Acetylene). Carborundum is a compound of silicon and carbon, hav- ing the composition SiC. It is made in the electric furnace by fusing sand (silicon dioxide, SiO 2 ), coke, saw.dust, and common salt. The essential chemical change is repre- sented thus : SiO 2 + 3C = SiC + 2 CO Silicon Dioxide Carbon Carborundum Carbon Monoxide Carborundum is silicon carbide (or carbon silicide). It is a crystallized solid, varying in color from white to emerald green and is sometimes iridescent. It is extremely hard, being harder than ruby and nearly as hard as diamond. Hence it is made into grinding wheels, whetstones, and polishing cloths. Over three million pounds were made at Niagara Falls in 1902, and the output is constantly increasing. Carborundum is a good conductor of heat. Its specific gravity is about three. Acids have no action upon it, but it is decomposed by fusing with potassium hydroxide and other alkalies. Carborundum is manufactured in a huge electric furnace, shown in Figure 18. It is an oblong box of bricks with permanent ends and loosely built sides. Each end is provided with a heavy metal plate. The wires for the electric current are attached to the outer ends of these plates, while the huge carbon electrodes fit into the inner ends, and project into the furnace. A cylinder of granulated coke makes an electrical connec- tion between the electrodes. In this furnace the rnixture is not heated by an electrical arc, but by the resistance of the carbon core to the pas- sage of the powerful current of electricity. The chemical change, as in the manufacture of calcium carbide, is due solely to heat. The current is passed through the mixture for about eight hours. When the opera- tion is over and the furnace is cool, the side walls are pulled down, and the carborundum is removed. The purest grade is found around the core. It is crushed, treated with sulphuric acid to remove the impurities, washed, dried, and graded according to the size of the particles. n8 Descriptive Chemistry. Artificial Graphite is formed in the manufacture of carborundum. It is also made by heating a certain grade of anthracite coal in an electric furnace. It is extensively used in making electrodes for electric furnaces. Over 800,000 Ib. were manufactured in 1902 at Niagara Falls. Graphite is a form of carbon (see Graphite). Light, Heat, Electricity, and Chemical Action. 119 ELECTRICITY AND CHEMICAL ACTION. The Relation between Electricity and Chemical Action has always been a fascinating subject. Volta constructed his voltaic pile about 1800. This was one of the first, per- haps the first, source of an electric current. In May, 1800, Nicholson and Carlisle decomposed water into hydrogen and oxygen by an electric current obtained from a thermo- pile. In the same year Cruikshank obtained lead and copper from solutions of their salts. And in 1807 Davy isolated the elements, sodium and potassium, by passing an electric current (obtained from a large battery) through fused caustic soda and caustic potash respectively. From that time until the present day, the relation between elec- tricity and chemical action has engaged the attention of chemists. And their labors have built up a branch of chemistry called electrochemistry, which has recently attained considerable com- mercial importance. The Voltaic (or Galvanic) Cell in its simplest form consists of two metals connected by a wire and dipped into a liquid which will interact with one of the metals (Fig. 19). Copper, zinc, and water containing sulphuric acid may be used as an illustration. When the connected metals are FlG I9 ._voltaic cell, put into the acid, the zinc slowly disappears and hydrogen bubbles appear on the copper. Further examination would show that the zinc and sulphuric acid interacted, forming zinc sulphate. The chemical change is the one already described under hydrogen, and may be represented thus : Zn + H 2 S0 4 H 2 + ZnS0 4 Zinc Sulphuric Acid Hydrogen Zinc Sulphate The connecting wire becomes electrified and exhibits the effects of an electric current, viz., it becomes warm, it makes a magnetic needle move, I2O Descriptive Chemistry. and a shower of sparks is produced if the wire is cut and one end is drawn down a file while the other is held firmly upon it. The source of the electric current is obviously the chemical action between the acid and zinc. The copper is necessary, otherwise the product of the chemical action would be merely heat. Carbon is often used in place of copper, and other liquids instead of sulphuric acid. The liquid chosen, how- ever, must be one that will interact with zinc or its substitute. Several cells joined together form an electric battery. For many years the battery was the chief source of the electric current. And it is now used, especially for ringing telephone, house, fire alarm, and signal bells, and in operating the telegraph. The dynamo is now widely used to generate powerful currents of electricity. Electrochemical Terms. Faraday (1791-1867) investi- gated electrochemistry about 1834, and introduced many terms in common use. He called the decomposing process electrolysis, and the decomposable liquid the electrolyte ; the wire by which the current entered he called the anode ; and that by which it escaped, the cathode. " Finally," he says, ^require a term to express those bodies which pass to the electrodes) I propose to distinguish such bodies by calling those anions which go to the anode of the decom- posing body ; and those passing to the cathode, cations ; and when I have occasion to speak of these together, I shall call them ions. Thus, chloride of lead is an electro- lyte, and when electrolyzed evolves the two ions, chlorine and lead, the former being an anion and the latter a cation." These terms are so used to-day, but they demand a broader definition. Electrolysis is the series of chemi- cal changes caused by the passage of an electric current through a dissolved or fused (i.e. melted) compound. The compound thus decomposed is an electrolyte. The metallic or carbon rods which conduct the current of electricity to and from the electrolyte are called the poles, or better, the electrodes. Electrodes are usually made of platinum, cop- Light, Heat, Electricity, and Chemical Action. 121 per, zinc, mercury, or hardened carbon ; they may have any shape rod, wire, sheet, plate, box, crucible ; and they may also be solid, liquid, or powder, as well as fixed or movable. The electrodes are connected by wires with the source of the electric current, and serve as "doors" to quote Faraday again for the current to flow into and out of the electrolyte and through the wire connecting the electrodes. We speak of a " current" of electricity and of electricity as " flowing," although we do not know the nature of electricity, nor do we mean really that it flows, like a river, only in one direction. It is customary to speak of the current as entering the electrolyte by the anode or positive electrode and leaving by the negative electrode or cathode. The anode is the electrode that is often con- sumed or worn away, either mechanically or chemically. But solids are often deposited upon the cathode, as will soon be described, ^ons are those parts of the decomposed electrolyte which are believed to be material carriers of elec- tricij^y Aw comes from a Greek word which means wander- ing or migrating. And a cation is that ion which moves down or along with the current of electricity to the cathode where it is separated, deposited, or modified; while an anion is that ion which moves upward or against the current to the anode, where it likewise appears in various forms. Anions are electro-negative ions, but cations are electro-positive ions. Metallic ions are cations ; hence metals are deposited at the negative electrode or cathode. Non-metallic ions are usually anions, therefore oxygen, chlorine, and their oxides and hydroxides appear at the anode. Hydrogen is electro-positive. In general, metals are electro-positive, and non-metals (except hydrogen) are electro-negative. Ions follow the law of electric attraction and repulsion, viz., ions with the same kind of electrification repel each 122 Descriptive Chemistry. FIG. 20. Electrolytic cell. A and C are the electrodes, R is the electrolyte, B or D is the battery or dynamo. other, and those with unlike kinds attract. Hence the electro-positive cations move toward the electro-negative cathode, and the electro-negative anions move toward the electro-positive anode. Ions are further described under lonization (see below). An elec- trolytic cell is the apparatus in which electrolysis takes place (Fig. 20). Its parts are analo- gous to the voltaic cell. There must be a containing Vessel, the two electrodes, and the electro- lyte. The vessel may have any desired shape, and is made of material which will resist the corrosive action of the electro- lyte or which will withstand a high temperature. Unlike the voltaic cell, the electrolytic cell generates no electric current ; it receives the current from a dynamo or a battery. Elec- trolysis is accomplished on a large scale in electrolytic cells. Illustrations of Electrolysis. Electrolysis may be simple, but it is usually very complex. Two illustrations will be given. When two platinum electrodes are put into melted zinc chloride and a current of electricity is passed, zinc is deposited at the cathode, and chlorine gas is liber- ated at the anode. This is a simple instance of electrolysis. But when an aqueous solution of sodium chloride is electro- lyzed, the action is different. Theoretically, the products should be sodium and chlorine, but they are hydrogen, sodium hydroxide, and chlorine. The sodium separated at the cathode immediately interacts with the water to form hydrogen and sodium hydroxide. Furthermore, unless the chlorine and sodium hydroxide are removed, they will interact to form compounds of chlorine, which vary in composition with the temperature, etc. Light, Heat, Electricity, and Chemical Action. 123 The Electrolysis of Water is more complex than is ordinarily sup- posed. Strictly speaking, it is the sulphuric acid, and not the water, that is electrolyzed. Perfectly pure water does not conduct electricity, and is consequently not decomposed by it. But since the same amount of sulphuric acid is always present, no matter how long the action con- tinues, it is customary to speak of the total change as the electrolysis of water. The hydrogen and oxygen gases, which collect at the cathode and anode respectively, are merely the end products of a series of changes. Small quantities of ozone and hydrogen dioxide are also formed. Faraday's Law. In his study of electrolysis, Faraday found that a measured quantity of electricity liberated different but definite amounts of the chemical elements. For example, the current which liberated i gm. of hydrogen also liberated 8 gm. of oxygen, 35.5 gm. of chlorine, 108 gm. of silver, 31.7 gm. of copper, and so on. These numbers are identical with the chemical equivalents of these . elements (compare Equivalents, Chapter IX) . Faraday called them electrochemical equiv- alents, to emphasize their chemical and electrical relationship. But the term electrochemical equivalent now means, however, the weight of an element deposited or liberated by a current of a certain arbitrary value (i ampere in I second) . For example, the electrochemical equivalent of hydrogen is 0.000010441 gm., of oxygen is 0.00008287, and sometimes 0.00016574, of copper is 0.0003294, and sometimes 0.0006588, of silver is 0.001118. This general relation is often stated as Faraday's Law, thus : When the same quantity of electricity acts upon different electrolytes, the ratio between the quantities of liberated products is the same as between their chemical equivalents. Faraday also showed that the amount of decomposition the chem- ical work, we might say is proportional to the total amount of elec- tricity used. It makes no difference whether the current is strong or weak, nor whether the time of its flow is long or short. A certain quantity of electricity will do so much chemical work no more and no less. Thus a given quantity of electricity passed through copper sulphate solution always deposits the same weight of copper at the cathode. These two principles of Faraday are at the foundation of all electrochemical industries. Their importance can hardly be over- estimated. 124 Descriptive Chemistry. Industrial Applications of Electrolysis. The earliest industrial application of electrolysis was in electrotyping and electroplating. These operations consist in depositing a thin film of metal upon a surface. They are fundamen- tally the same, though copper is the only metal used for producing electrotypes. Electrotypes are exact repro- ductions of the original objects. The process of electro- typing is substantially as follows : the page of type, or the woodcut, is first reproduced in wax or plaster. This exact impression is next covered with powdered graphite to make it conduct electricity. The coated mold is then suspended as the cathode in an acid solution of copper sulphate ; the anode is a plate or bar of copper. When the current is passed, electrolysis occurs ; copper is dis- solved from the anode and deposited upon the mold in a film of any desired thickness. The exact copper copy is stripped from the mold, backed with metal and mounted on a wooden block, and used instead of the type or woodcut itself. By this process exact copies of expensive wood engravings can be cheaply reproduced, and type can be saved from the wear and tear of printing. Most books, magazines, and newspapers are now printed from electro- types. The process of electroplating differs from elec- trotyping in only one essential, viz., in electroplating, the deposited film is not removed from the object. The object to be plated is carefully cleaned and made the cathode ; the anode is a bar or plate of the metal to be deposited. When the current passes through the system, the metal is firmly deposited upon the object. The electrolysis would take place, of course, if any anode were present ; but anodes of the metal to be deposited are usually used to prevent the solution or " bath " from weakening. They accom- plish the purpose by replenishing the solution with metal Light, Heat, Electricity, and Chemical Action. 125 as fast as it is removed and deposited upon the cathode. Silver, nickel, and gold are the usual metals used in electroplating (see these metals). Electroplating and electrotyping have been done since about 1840. It is only within the last ten or fifteen years, however, that the electric current has been profitably applied in many industries. But during this time the development of electrochemistry has been very marked. The largest of these industries is the refining of copper. The process is similar to that described under electro- typing. Other metals, such as gold, silver, and lead, are extracted from their ores and purified by electricity, though the older processes are still used. All the aluminium, mag- nesium, and sodium of commerce are now manufactured by passing an electric current through their fused compounds. Nearly all the domestic potassium chlorate and much of the caustic soda are made by electricity. The same is true of barium compounds and many other chemicals. These electrochemical processes will be fully discussed in the appropriate places. The Theory of Electrolysis. Many theories have been proposed to explain electrolysis. According to the theory now generally held, electrolysis is not the splitting or tear- ing apart of molecules by the electric current. It is the carrying of electricity from one electrode to the other by ions. Dissolved or fused compounds are more or less dis- sociated into ions before the current of electricity is intro- duced, /and the current flows simply because the ions are there to carry it. Since these ions are charged with elec- tricity, the dissociation is called electrolytic dissociation or ionization. Ions are not atoms, but electrically charged atoms or groups of atoms. Thus, when sodium chloride is dissolved in water, much of the salt dissociates into the 126 Descriptive Chemistry. ions, sodium and chlorine ; the sodium ions are charged positively, and the chlorine atoms negatively. Now, when an electric current is passed into the solution, the ions move toward their proper electrodes, carrying the electric charges with them. In brief, the current sorts the ions, which in turn migrate with their charges. When the ions reach their respective electrodes, they give up their electric charges and assume their normal conditions. Thus, the positive sodium ions give up their charges at the negative electrode, or cathode, and become sodium atoms. The latter interact with water to form hydrogen and sodium hydroxide. Similarly, the negative chlorine ions give up their charges at the positive electrode, or anode, and become neutral atoms, which at once unite to form chlo- rine molecules. Electrolysis and Solution. According to the above theory, the properties of many water solutions are closely related to the phenomena of electrolysis. For many years it was believed that a dissolved substance was distributed unchanged throughout the solvent. It was also believed that certain dissolved substances combined in part with the water a view held to-day. The first real step toward a settlement of the problem was taken when the electri- cal conductivity of solutions was compared. Experiments show that the electrical conductivity of solutions varies between wide limits. Water itself is practically a non- conductor, a sugar solution is a very poor conductor, while solutions of most acids, bases, and salts are excellent con- ductors. Water solutions, therefore, are of two kinds : (i) those which conduct electricity, and (2) those which do not, or only very slightly. But we have already seen that the first class consists of electrolytes. Hence, two things are believed about water solutions: (i) that when Light, Heat, Electricity, and Chemical Action. 127 acids, bases, and salts are dissolved in water, they are dis- sociated into ions, and (2) that when sugar and similar substances are dissolved in water they dissociate very slightly or not at all. The amount of dissociation depends largely upon the relative amounts of solute and solvent, i.e. upon the dilution of the solution. The dissociation is slight in concentrated solutions, but increases as the dilu- tion increases. Not all acids, bases, and salts dissociate to the same degree. The percentage of dissociation of some of these compounds in solutions of a certain strength and at the same temperature (i 8 C.) is given in the following TABLE OF IGNIZATION. SUBSTANCE. . PER CENT OF IONIZATION. Hydrochloric acid 78 / 82 Potassium chloride 7C Potassium nitrate / j 64. Potassium hydroxide 77 Sodium hydroxide 7? Numerous facts support the theory of ionization. (i) Varying electrical conductivity has already been mentioned. (2) It has long been known that solutions boil at a higher temperature and freeze at a lower temperature than pure water. A fresh-water river, for example, freezes before the ocean, and water containing considerable mineral mat- ter boils at a higher temperature than pure drinking water. It is generally true that a dissolved substance raises the boiling point and loivers the freezing point of a given solu- tion. Now, when weights of substances proportional to 128 Descriptive Chemistry. their molecular weights are dissolved in the same volume of water, the boiling point of each solution is raised the same number of degrees and the freezing point is lowered the same number of degrees. These facts are now applied experimentally to determine molecular weights. In many cases the molecular weights thus found agree with the values obtained by other methods. Thus, if X is the depression produced by a one per cent solution of sugar, and Y the depression produced by a one per cent solution of urea, the following proportion may be written, because the depressions of the freezing points are inversely propor- tional to the molecular weights Y: X : : mol. wt. of sugar : mol. wt. of urea. The molecular weight of sugar is known to be 342, and from the proportion the molecular weight of urea is 60, which agrees with that found by other methods. This method is applicable to many compounds and is helpful in deciding whether a molecular weight is a given number or its multiple. There is a marked disagreement to this rule, however, in the case of solutions of acids, bases, and salts. That is, electrolytes are exceptions. In some instances the molecular weight is only half that found by other methods. Thus, the molecular weight of sodium chloride was found to be about 30, instead of 58.5 the correct molecular weight. Hence, it is believed that the solutions of acids, bases, and salts contain ions which act like molecules in their effect upon the freezing and boiling points of solutions. The behavior of acids, bases, and salts in solution led the Swedish chemist Arrhenius, in 1887, to extend the ideas of Faraday and to propose the present theory of solution. Light, Heat, Electricity, and Chemical Action. 129 Application of the Theory of lonization. Many ob- scure facts of chemistry become intelligible when inter- preted by the theory of ionization. (i) Ordinary tests are tests for ions. For example, all chlorides in solution have the same test. That is, they all interact with silver nitrate in solution, because all have chlorine ions in the solution. Similarly, all soluble sulphates interact with barium chlo- ride in solution, because all sulphates have SO 4 ions in the solution. Both silver chloride and barium sulphate are insoluble, and are removed from the solution as precipi- tates. A complete illustration will make this fact clearer. The silver nitrate and sodium chloride solutions before mixing consist largely of the ions of silver, NO 3 -group, sodium, and chlorine. When mixed, the ions of silver and chlorine unite to form silver chloride, which is in- soluble and hence not ionized ; the solution still contains ions of sodium and of the NO 3 -group. On the other hand, if solutions of potassium chlorate and silver nitrate are mixed, no silver chloride is formed, because no chlo- rine ions are available. Potassium chlorate dissociates into ions of potassium and C1O 3 . Equations are often used to express ionization. Thus, the ionic equation for the interaction of sodium chloride and silver nitrate is Na + Cl 4 Ag + NO 8 = AgCl 4- Na 4 NO 3 . (2) lonization explains the General Properties of Acids, Bases, and Salts. Acids in solution turn litmus red, be- + cause their solutions contain hydrogen ions (H). Simi- larly, bases turn litmus blue, because their solutions contain hydroxyl ions (OH). But solutions of neutral salts con- tain neither hydrogen nor hydroxyl ions, hence they do not affect litmus. The above principles can be readily i jo Descriptive Chemistry. extended to cover acid and basic salts. The other general properties of acids and bases are believed to be due to the above causes. (3) Neutralization, interpreted by the ionic theory, is fundamentally the union of hydrogen and hy- droxyl ions to form molecules of water. Suppose hydro- chloric acid and potassium hydroxide are mixed. The solution at first contains the hydrogen, chlorine, potassium, and hydroxyl all as ions. But the hydrogen and hy- droxyl immediately unite to form water, leaving the po- tassium and chlorine ions in the solution. This solution is thus rendered neutral by the removal of the hydrogen ion its acid constituent and of the hydroxyl ion its basic constituent. The ionic equation expressing the neutralization of potassium hydroxide by hydrochloric acid is K + OH + H + Cl = K + Cl + H 2 O. The potassium and chlorine ions remain free and un- combined until the solution is evaporated. As the con- centration increases, the ions unite until nothing remains except* the neutral salt potassium chloride. Neutralization, therefore, as interpreted by the ionic theory, is essen- tially a union of hydroxyl and hydrogen ions. This view is supported by much experimental evidence. For example, the heat of neutraliza- tion produced by the interaction of equivalent quantities of strong acids and bases is approximately the same. EXERCISES. 1. What transformations of energy accompany chemical action? Illustrate your answer. 2. State and illustrate the law of the conservation of energy. 3. Discuss the relation of light to chemical action. Give popular and scientific illustrations of (a} the production of chemical action by light, and (#) production of light by chemical action. 4. Define and illustrate (a} calorie, (6) thermal equation, (c} heat of formation, (d} exothermic, (e) heat of decomposition, (/) endothermic. Light, Heat, Electricity, and Chemical Action. 131 5. Give several illustrations of the production of (a} heat by chemi- cal action, and (b) vice versa. 6. When an electric spark is passed through a mixture of two vol- umes of hydrogen and one volume of oxygen, what is the result? Is it due directly to electricity or to heat? 7. Define and illustrate kindling temperature. 8. Name several sources of heat. How may electricity be used as a source of heat ? 9. Describe Moissan's electric furnace. Why is it so efficient? Is its effect thermal or electrical ? State some results produced by Moissan with this furnace. Has the electric furnace any industrial use ? Where ? 10. What is calcium carbide? How is it made? State the equation for the reaction. What are its properties ? For what is it used? 11. What is carborundum? How is it made? State the equation for the reaction. What are its properties and uses? 12. What is artificial graphite? How is it made? For what is it used? 13. Give several illustrations of the production of (a) electricity by chemical action, and (b) vice versa. 14. State briefly the first chemical changes which were produced by electricity. 15. Describe a simple voltaic cell. Why is it so called? What is the source of the electric current manifested by the cell? What is an electric battery ? For what is it used ? 1 6. Define and illustrate (a) electrolysis, (b) electrolyte, (c} elec- trode, (W) anode, (e) cathode, (/) ions, (g) anion, (h} cation, (*) posi- tive electrode, (/) negative electrode, () ionization. 17. Where are (a) anions and (b) cations liberated? 1 8. Describe an electrolytic cell. How does it differ from a voltaic cell ? For what is it used ? 19. Describe the electrolysis of (a) zinc chloride, (b} sodium chloride, (c) water. 20. State and illustrate Faraday's law. 2 1 . Give a brief account of Faraday's contribution to electrochemistry. 22. Describe the process of (a) electrotyping and (b} electroplating. 23. State some industrial applications of the electric current. 24. What is the theory of electrolysis? What is the present theory of solution in water? What is the theory called? Why? What facts support it? 25. Define and illustrate an ionic equation. Descriptive Chemistry. PROBLEMS. 1. Calculate the percentage composition of (a) water, () magnetic oxide of iron (Fe 3 O 4 ), (c} crystallized sodium carbonate (Na.,CO 3 . ioH 2 O). 2. If a certain current of electricity deposited 31.7 gm. of copper, how much (a) silver, (&) aluminium, and (c) magnesium would it deposit ? 3. If a certain current of electricity deposited 2 kg. of copper, how much silver would it deposit? 4. How much calcium carbide can be made (theoretically) from a ton of lime? (Equation is 3 C + CaO = CaC 2 + CO or 36 + 56 = 64 + 28.) 5. How much carborundum can be made (theoretically) from a ton of sand (SiO 2 ) ? (Equation is SiO 2 + 3 C = SiC + 2 CO or 60 + 36 = 40 + 56.) 6. Calculate the percentage composition of (#) carborundum and (&) calcium carbide. CHAPTER XL CHLORINE AND HYDROCHLORIC ACID. CHLORINE is an important element, and its compounds are useful, especially hydrochloric acid, sodium chloride, and bleaching powder. Occurrence. Free chlorine is never found in nature, because it combines so readily with other elements. But in combination it is widely distributed, since it is one of the components of common salt, or sodium chloride. Many compounds of chlorine with potassium, magnesium, and calcium are found in the deposits at Stassfurt in Ger- many (see these metals). The salts found in sea water contain about 2 per cent, and the earth's crust contains about o.oi per cent of chlorine. Silver chloride "horn" silver is mined as an ore in the United States and Mexico. Preparation. Chlorine is prepared in the laboratory by heating a mixture of manganese dioxide and hydro- chloric acid. This method was used by Scheele, who discovered the gas in 1774. The equation for the prepa- ration of chlorine is MnO 2 + 4HC1 = C1 2 + MnCl 2 + 2 H 2 O Manganese Hydrochloric Chlorine Manganese Water Dioxide Acid Bichloride V This is an oxidizing process, since the hydrogen of the hydrochloric acid is oxidized to water, although only part of the chlorine of the acid is obtained free. 134 Descriptive Chemistry. Sometimes chlorine is prepared in the laboratory by heating a mixture of manganese dioxide, sodium chloride, and sulphuric acid. This method is substantially the same as the other, since a mixture of sulphuric acid and sodium chloride yields hydrochloric acid. The simplest equation for this method of preparing chlorine is 2 H 2 SO 4 + 2 NaCl + MnO 2 = C1 2 + Na,SO 4 + MnSO 4 -f 2 H 2 O Sulphuric Sodium Manganese Chlorine Sodium Manganese Water Acid Chloride Dioxide Sulphate Sulphate Other oxidizing substances besides manganese dioxide may be used, such as potassium chlorate (KC1O 3 ), potassium dichromate (K 2 Cr 2 O 7 ), and red lead (Pb 3 O 4 ). Chlorine is manufactured by several processes, all of which involve the same principle as the laboratory method. In the Deacon process, hydrochloric acid is oxidized by oxygen ob- tained from the atmosphere. A mixture of hydrochloric acid gas and air is heated to 500 C. and passed through iron tubes containing balls of clay or pieces of brick previously saturated with copper chloride. A series of complex reactions occurs which are not well understood. It is supposed that the copper chloride facilitates the formation of chlorine by continuously giving and taking this gas. The essential chemical change, however, is the oxidation of the hydrochloric acid, and it may be represented by the equation 2HC1 -I- O C1 2 -f H 2 O Hydrochloric Acid Oxygen Chlorine Water In the Weldon process, an impure native manganese dioxide, known as pyrolusite, is treated with hydrochloric acid in large earthenware retorts or stone tanks heated by hot water or steam. When no more chlorine is liberated, the residue is mainly manganese dichloride. This " still- liquor" was formerly thrown away, but by the Weldon process it is changed into manganese compounds, which are used to prepare more chlorine (see Manganese Dioxide). Chlorine is also prepared on a large scale by the electrolytic process. Sodium chloride is decomposed by electricity in properly constructed cells, and the chlorine which is liberated at the anode is conducted off through pipes to the bleaching powder factory. Sodium hydroxide is produced at the same time, and the process will be described under this compound. Chlorine and Hydrochloric Acid. 135 Properties. Chlorine is a greenish yellow gas. Its color suggested the name chlorine (from the Greek word chloros, meaning greenish yellow), which was given to it by Davy about 1810. It has a disagreeable, suffocating odor, which is very penetrating. If breathed, it irritates the sensitive lining of the nose and throat, and a large quantity would doubtless cause death. It is heavier than the other elementary gases, and is about 2.5 times heavier than air. Hence it is easily collected by downward dis- placement, i.e. by allowing it to fall to the bottom of a bottle and thus fill the latter by displacing the air. A liter of dry chlorine at o C. and 760 mm. weighs 3.18 gm. Water dissolves chlorine. The solution is yellowish, smells strongly of chlorine, and is frequently used in the laboratory as a substitute for the gas. Chlorine water, as the solution is called, is unstable even under ordinary con- ditions, and must be kept in the dark. If the solution is placed in the sunlight, oxygen is soon liberated and hydro- chloric acid is formed. Intermediate changes doubtless occur ; but the simplest equation for the essential change is- Hp + C1 2 = 2HC1 + O Water Chlorine Hydrochloric Acid Oxygen Chlorine is much less soluble in a solution of sodium chloride, over which it is sometimes collected. It attacks mercury and cannot be col- lected over this liquid. Chlorine does not burn in the air, but many substances burn in chlorine. The metals antimony and arsenic, when sprinkled into chlorine, suddenly burst into flame, while phosphorus melts at first and finally burns with a feeble flame. If sodium, iron powder, brass wire, or other metals are heated and then put into chlorine, they burn ; the' sodium and iron produce a dazzling light and the brass 136 Descriptive Chemistry. glows and emits dense fumes of whitish smoke. Chlorine combines readily with hydrogen. Hence, a jet of burning hydrogen when lowered into chlorine continues to burn, forming hydrochloric acid gas, which appears as a white cloud. The simplest equation for this change is H + Cl HC1 Hydrogen Chlorine Hydrochloric Acid The attraction between chlorine and hydrogen is so great that many compounds of hydrogen are decomposed by chlorine. Thus, compounds containing hydrogen and carbon, such as illuminating gas, paraffin wax, and wood, burn in chlorine with a smoky flame. Chlorine does not combine directly with carbon, hence the flame consists largely of very fine particles of solid carbon. Similarly, a piece of glowing charcoal is extinguished by chlorine. If filter paper is saturated with warm turpentine (a compound of hydrogen and carbon) and put into a bottle of chlorine, a flame accompanied by a dense cloud of black smoke bursts from the bottle ; the chlorine withdraws the hydro- gen to form hydrochloric acid, while the carbon is left free. The power to bleach is the most striking and useful property of chlorine. This property depends upon the fact, already mentioned, that chlorine withdraws hydrogen and liberates free oxygen ; the latter then decomposes the coloring matter in the cloth or other material. Dry chlorine does not bleach. If an envelope on which the postmark, or a lead pencil mark, is still visible is placed in moist chlorine, these marks will not be bleached be- cause they are largely carbon ; but the writing ink, which is mainly a compound of hydrogen, carbon, and iron, will disappear. Litmus paper and calico are both bleached by moist chlorine. Chlorine and Hydrochloric Acid. 137 Bleaching Powder is the source of the chlorine used in the bleaching industries. It is sometimes called "bleach," or " chloride of lime." It is a yellowish white substance having a peculiar odor, which resembles that of chlorine. When dry, it is a powder, but on exposure to the air, it absorbs water and carbon dioxide, becomes lumpy and pasty, and loses some of its chlorine. Acids like sulphuric and hydrochloric acid liberate from bleaching powder its " available chlorine," which varies from 30 to 38 per cent in good qualities. The equations for the interaction of acids and bleaching powder are usually written thus CaOCl 2 + H 2 SO 4 = Cl a + CaSO 4 + H 2 O Bleaching Powder Sulphuric Acid Calcium Sulphate CaOCl 2 + 2 HC1 = C1 2 -f CaCl 2 + H 2 O Hydrochloric Acid Calcium Chloride The composition of bleaching powder has been much discussed. The most reliable authority gives it the formula CaOCL,. When dis- solved in water, bleaching powder forms calcium hypochlorite (CaO 2 Cl 2 ) and calcium chloride (CaCl.,). Bleaching Powder is manufactured by the action of chlorine gas on lime. Lime (calcium oxide, CaO) is carefully slaked with water to form calcium hydroxide (Ca(OH). 2 ). This powder is sifted into a large absorption chamber made of iron, lead, or tarred brick until the floor is covered with a layer three or four inches deep. The chlorine enters at the top and settles slowly to the floor, where it is absorbed by the lime. The simplest equation for the formation of bleaching powder might be written Ca(OH) 2 + C1 2 CaOCl 2 + H 2 O Calcium Hydroxide Chlorine Bleaching Powder Water Bleaching. Immense quantities of bleaching powder are used to whiten cotton and linen goods and paper pulp. The pieces of cotton cloth as they come from the mill are 138 Descriptive Chemistry. sewed end to end in strips, which are stamped at the extreme ends with some indelible mark to distinguish each owner's cloth. These strips, which are often several miles long, are drawn by machinery into and out of numerous vats of liquors and water, between rollers, and through machines, until they are snow-white and ready to be finished (i.e. starched and ironed) or dyed. The whole operation requires three or four days. The preliminary treatment consists in singeing off the downy pile and loose threads by drawing the cloth over hot copper plates or through a series of gas flames. The object of the remaining operations is threefold, (i) to wash out mechanical impurities, the fatty and resin- ous matter, and the excess of the different chemicals, (2) to remove matter insoluble in water, and (3) to oxidize the coloring matter by chlorine. The details of the process differ with the texture of the cloth and with its ultimate use. The threefold object above mentioned involves successively "liming," "souring," "chemicking," and "souring, 11 interspersed with frequent washing. The "liming 11 consists in boiling the cloth in a large kier or vat with lime, the "souring 11 in wetting it with weak sulphuric or hydrochloric acid, and the " chemicking ?1 in im- pregnating it with a weak solution of bleaching powder. Often the cloth is boiled at a certain stage with resin and sodium carbonate. The ^liming 11 removes the resinous and the fatty matter, the first "souring 11 neutralizes traces of lime, and the second, which follows the "chem- icking, 11 liberates the chlorine in the fiber of the cloth. Frequent washing is absolutely necessary to remove the impure products of the chemical changes as well as the excess of lime and other alkali, acid, and chlo- rine. Should these be left, the cloth would be unevenly bleached and its fiber would be weak. The cloth is finally treated with an antichlor, such as sodium hyposulphite, which removes the last traces of chlorine. Bleaching is chemically an oxidizing process. The oxygen when it is liberated from water by chlorine is said to be in the nascent state. This means that the gas is exceedingly active, because it is not only uncombined, but just ready to unite with those elements for which it has great affinity. Hence this nascent oxygen literally tears Chlorine and Hydrochloric Acid. 139 down complex colored substances and changes them into colorless compounds. The nascent state is aptly illustrated by bleaching because both the chlorine and the oxygen are in this active chemical condition. Chlorine Hydrate is formed by cooling chlorine water or by passing chlorine into ice water. It is a yellowish, crystalline solid, and in the air it decomposes quickly into chlorine and water. Its composition corresponds to the formula C1 2 10 H 2 O. Liquid Chlorine was first prepared by Faraday in 1823. A little chlorine hydrate was inclosed in one arm of a bent tube (Fig. 21), which was then sealed. By gently heating the tube, the chlorine hy- drate was decomposed into chlorine and water, but the chlorine, being unable to escape, was condensed to a liquid by the pressure inside the tube. The liquefaction is more easily accom- plished if one end is kept cold during the experiment. FlG - 21.- Bent tube for . . ,. ,, . .the liquefaction of chlo- At the ordinary pressure, chlorine gas be- r j ne comes liquefied, if its temperature is 34 C, while at a pressure of six atmospheres the temperature need be only o C. Liquid chlorine has a bright yellow color. It is a commercial article, and is stored and shipped in steel cylinders lined with lead. It is used in the laboratory to prepare chlorides, and industrially to extract gold. Solid chlorine has been obtained as a yellow crystalline mass by cooling the liquid to 102 C. Uses of Chlorine. Chlorine is used directly to prepare some of its compounds, the most important being bleaching powder. The latter is often used as a deodorizer and dis- infectant, since the liberated chlorine destroys putrefying matter by acting on it as on coloring matter. A solution of potassium hypochlorite (Javelle's water) or sodium hy- pochlorite (Labarraque's solution) is often used to remove fruit stains from cotton and linen goods. Chlorides are formed when chlorine combines with other elements, and they are in general stable compounds. 140 Descriptive Chemistry. The simplest equations illustrating the combination of chlorine with metals and other elements are Na + Cl = NaCl Sodium Chlorine Sodium Chloride Sb + 3d = SbCl 3 Antimony Antimony Trichloride Cu + C1 2 = CuCl 2 Copper Copper Chloride P + 3 C1 = PC1 3 Phosphorus Phosphorus Trichloride H + Cl = HC1 Hydrogen Hydrochloric Acid Chlorides form an important class of compounds and they will be considered under the elements with which the chlorine combines. (See also Chlorides below.) HYDROCHLORIC ACID. Hydrochloric Acid is the most useful compound of chlorine. It is a gas, very soluble in water^ This solution has long been known as muriatic acid (from the Latin word muria, meaning brine). The term hydrochloric acid includes both the gas and its solution, but the solution is usually meant. The early chemists called the gas " spirit of salt." Priestley, who first prepared, collected, and studied the gas, called it " marine acid air." Both expressions emphasize its relation to salt (sodium chloride). Occurrence. The gas occurs free in volcanic gases. The solution is one constituent of the gastric juice of the stomach. Chlorides, which are salts of hydrochloric acid, are abundant in the earth's crust. Preparation. The gas is prepared in the laboratory by the method devised by Glauber in the seventeenth cen- Chlorine and Hydrochloric Acid. 141 tury, viz., by heating sulphuric acid and sodium chloride. If the mixture is gently heated, the chemical change is represented thus Nad + H 2 S0 4 = HC1 + HNaSO 4 Sodium Sulphuric Hydrochloric Acid Sodium Chloride Acid Acid Sulphate But at a high temperature the equation for the reaction 2 NaCl + H 2 SO 4 - 2 HC1 + Na 2 SO 4 In either case the gas is readily produced. It may be collected over mercury or, more easily, by downward dis- placement. The solution is prepared by passing the gas into water. That sodium sulphate is the other product of the chemical change at a high temperature may be shown by testing the heated residue as follows : (a) Dissolve a portion in water and add a few drops of barium chloride solution ; the immediate formation of the white, insoluble barium sulphate shows that the residue from the experiment must be a sulphate, (b} Burn a little of the residue on a platinum wire or piece of porcelain held in the Bunsen flame ; the intense yellow color immediately imparted to the flame shows that the residue contains sodium, (c) Hence the compound must be sodium sulphate. Commercial Hydrochloric Acid is manufactured in enor- mous quantities by the method used in the laboratory. A mixture of salt and sulphuric acid is moderately heated in a large hemispherical cast-iron pan, and the gas passes through an earthenware pipe into an absorbing tower ; the fused mass of acid sodium sulphate and salt is then sub- jected to a higher temperature, and the liberated gas passes by another pipe into the absorbing tower. These towers are tall and filled with coke or pieces of brick over which water trickles ; as the hydrochloric acid gas passes up the tower, it is absorbed by the descending water, and flows 142 Descriptive Chemistry. out at the bottom of the tower as concentrated acid. The gas is usually cooled before it enters the towers. Some- times the gas passes through huge earthenware jars be- fore entering the towers. In these jars the gas and water are caused to flow constantly in opposite directions, thus insuring complete absorption. Hydrochloric acid gas is a by-product in the manufacture of sodium carbonate by the Leblanc process. The gas was formerly allowed to escape into the atmosphere, but since it destroyed vegetation and be- came a nuisance in other ways, a law was passed forbidding the manu- facturers to let it escape. Hence it became necessary to absorb the gas in water. The hydrochloric acid, which was once regarded as a waste product, is now the main source of profit, since competition has reduced the price of sodium carbonate (see Sodium Carbonate). Properties. Hydrochloric acid gas is colorless and transparent. When it escapes into moist air, it forms fumes which are really minute drops of a solution of the gas in the moisture of the air. It has a choking, sharp, pungent odor. The gas does not burn nor support com- bustion. It is about 1.25 times heavier than air, and may therefore be collected by downward displacement. One liter at oC. and 760 mm. weighs 1.61 gm. The gas can be liquefied at ioC. and 40 atmospheres pressure; while at i6C, the pressure need be only 20 atmospheres. The extreme solubility of hydrochloric acid gas in water is one of its most striking properties. One liter of water will dissolve about 500 1. of gas, if both are at o C. and 760 mm. At the ordinary temperature about 450 1. of gas dissolve in i 1. of water, and as the temperature rises the solubility decreases. The solution is the familiar hydrochloric acid. The gas readily escapes, hence the acid forms fumes when exposed to air. Pure hydrochloric acid is a colorless liquid. The commercial acid has a yellow color, usually due to iron Chlorine and Hydrochloric Acid. 143 compounds, but sometimes to organic matter or to dissolved chlorine. It also contains other impurities. Like most acids, it reddens blue litmus, and gives up its hydrogen when added to metals. The strongest acid contains about 42 per cent (by weight) of the gas, and its specific gravity is i .2. When the strong acid is heated, the gas is evolved until the solution contains about 20 per cent of the acid, and then the liquid boils at i ioC. without further change. The dilute acid, on the other hand, loses water until the same conditions prevail. Composition of Hydrochloric Acid Gas. In 1810, Davy showed that hydrochloric acid gas (which had been regarded as an oxygen compound) contained only chlorine and hydrogen. Many facts lead us to conclude that hydrochloric acid gas is composed of hydrogen and chlorine in such a ratio that its composition is represented by the for- mula HC1. (i) Hydrogen burns in chlorine, and the only product is hydrochloric acid gas. (2) When hydrochloric acid is decomposed by an electric current, equal volumes of hydrogen and chlorine are evolved. (3) When a mixture of equal volumes of hydrogen and chlorine is exposed to the direct sunlight or to the action of an electric spark, the gases combine with an explosion, and hydrochloric acid gas is formed with no residue. Furthermore, the volume of the resulting gas equals the sum of the volumes of hydrogen and chlorine used. (4) When a given volume of dry hydrochloric acid gas is treated with sodium amalgam, the chlorine is withdrawn by the sodium in the amal- gam, and a volume of hydrogen remains which is half the original vol- ume. (5) No derivative of hydrochloric acid is known which contains less hydrogen, or less chlorine in a molecule. (6) The ratio by weight in which hydrogen and chlorine combine is 1:35.45. Hence, the lowest molecular weight of hydrochloric acid is 36.45, a number which has been verified by several different methods. Uses of Hydrochloric Acid. Vast quantities are used to prepare chlorine for the manufacture of bleaching pow- der. Various chlorides are prepared from it, and it is one of the common acids used in chemical laboratories. Chlorides are formed by the direct addition of chlorine to metals, as we have seen. They are also formed when 144 Descriptive Chemistry. metals, their oxides, or hydroxides are added to hydro- chloric acid. The following equations illustrate this gen- eral fact : Zn + 2 HC1 = ZnCl 2 + H 2 Zinc Zinc Chloride ZnO + 2 HC1 = ZnCl 2 + H 2 O Zinc Oxide Zinc Chloride Zn(OH) 2 + 2 HC1 - ZnCl 2 + 2 H 2 O Zinc Hydroxide Zinc Chloride They are also formed by adding other salts to hydro- chloric acid. Molecules of chlorides may contain several atoms of chlorine. Occasionally the- name of the compound indicates this fact, e.g. manga- nese dichloride (MnQ 2 ), antimony trichloride (SbCl ;5 ), phosphorus trichloride and pentachloride (PC1 3 and PCI-)- If a metal forms two chlorides, the two are distinguished .by modifying the name of the metal. The one containing the smaller proportion of chlorine ends in -ous, the one containing the larger ends in -ic. Thus, mercurous chlo- ride is HgCl, but HgCl 2 is mercuric chloride. Similarly, we have fer- rous chloride, FeCl 2 , and ferric chloride, FeCl 3 . The Test for Hydrochloric Acid and Chlorides. Most chlorides are soluble in water. Those of lead, silver, and mercury (-ous) are not. If silver nitrate is added to hydro- chloric acid, or to the solution of a chloride, a white, curdy precipitate of silver chloride is formed, which (a) is insol- uble in nitric acid, but soluble in warm ammonium hydrox- ide, and () turns purple in the sunlight. The invariable formation of silver chloride is the test for hydrochloric acid and soluble chlorides. Hydrochloric acid gas also forms dense white clouds of ammonium chloride in the presence of ammonia gas. Chlorine and Hydrochloric Acid. 145 Miscellaneous. The acids of chlorine are tabulated under ACIDS. The compounds of chlorine with sodium, potassium, magnesium, and calcium are described under these metals. Aqua regia, of which chlorine is one constituent, is discussed in Chapter XII. EXERCISES. 1. What is the symbol of chlorine ? What useful compounds con- tain this element ? 2. How is chlorine prepared in the laboratory ? Give one equation for its preparation. Describe Deacon^ process for manufacturing chlorine. 3. Who discovered chlorine ? Who named it, when, and why ? 4. Summarize the physical properties of chlorine. How can it be quickly distinguished from the gases previously studied ? 5. Summarize the chemical properties of chlorine. Compare it with oxygen. Describe fully its action with hydrogen. 6. Define (a) downward displacement, (b} available chlorine, (V) antichlor. 7. Develop the topics : (a) nascent state, (<$) chlorine water, (V) chlo- rine hydrate, (ti ) liquid chlorine, (e) chlorine is an oxidizing agent. 8. What is bleaching powder ? How is it made ? What are its chief properties ? Describe the operation of bleaching. What is the chemistry of bleaching ? 9. What is (a) "bleach," () muriatic acid, (V) chloride of lime, (d) "salt, 11 (e) "lime," (/) commercial hydrochloric acid ? 10. What are chlorides ? Name five. How can they be formed ? Give the formula of sodium chloride. Why cannot chlorine be collected over mercury ? 11. What is hydrochloric acid ? How is it prepared in the labora- tory ? Give the equations for its preparation. How is it prepared industrially ? 12. Summarize the chief properties of hydrochloric acid gas. Of the acid, as the term is usually used. What happens when hydrochloric acid is boiled ? 13. What is the evidence that the formula of hydrochloric acid gas is HC1 ? 14. For what is hydrochloric acid used ? State the test for hydro- chloric acid and soluble chlorides. 146 Descriptive Chemistry. 15. Give a brief account of Faraday's work on chlorine. Of Davy's work. 1 6. Why is chlorine never found free ? PROBLEMS. 1. One equation for the preparation of chlorine is 4HC1 + MnO 2 = C1 2 + MnCI 2 + 2H 2 O 146 + 87 =71 + 126 + 36 (0) How many grams of chlorine can be made from 247 gm. of man- ganese dioxide ? () Name all the products. 2. How much sodium chloride is needed to prepare a kilogram of hydrochloric acid gas ? 3. How many grams of manganese dioxide are necessary to ^prepare 100 gm. of chlorine from hydrochloric acid. 4. A bottle of chlorine water was exposed to the sunlight until all the chlorine disappeared, (a) What two products were formed ? (^) Write the equation for the reaction. (c} What weight of chlorine gas is necessary to form 20 gm. of the gaseous product ? (d) What volume of chlorine is necessary to form 20 gm. of the other product ? 5. Calculate the percentage composition of (a) hydrochloric acid gas, (b) sodium chloride, (c) silver chloride (AgCl), (d) potassium chloride (KC1). CHAPTER XII. COMPOUNDS OF NITROGEN. THE most important compounds of nitrogen are am- monia (NH 3 ), nitric acid (HNO 3 ), and compounds related to them. Many animal and vegetable substances essential to life are compounds of nitrogen. AMMONIA. The term ammonia includes both the gas and its solu- tion in water, though the latter is more accurately called ammonium hydroxide. Formation of Ammonia. When vegetable and animal matter containing nitrogen decays, the nitrogen and hydro- gen are liberated in combination, as ammonia. The odor of ammonia is often noticed near stables. If animal sub- stances containing nitrogen are heated, ammonia is given off. The old custom of preparing ammonia by heating horns and hoofs in a closed vessel, i.e. by dry distillation, gave rise to the term "spirits of hartshorn." Soft coal contains compounds of nitrogen and of hydrogen, and when the coal is heated to make illuminating gas, one of the prod- ucts is ammonia. Preparation. Ammonia gas is prepared in the labora- tory by heating ammonium chloride with an alkali, usually slaked lime. The reaction may be represented thus H7 148 Descriptive Chemistry. 2NH 4 C1 + Ca(OH) 2 = 2NH 3 + CaCl 2 + 2 H 2 O Ammonium Slaked Ammonia Calcium Chloride Lime Gas Chloride 107 + 74 =34 + 111 + 36 The gas is usually collected by upward displacement, i.e. by allowing the gas to flow upward into a bottle and dis- place the air. The solution is prepared by conducting the gas into water. The main source of the ammonia of commerce is the ammoniacal liquor or gas liquor of the gas works. The gases which come from the retorts in which the coal is heated are passed into water, which absorbs the ammonia and some other gases. This impure gas liquor is treated with lime to liberate the ammonia, which is absorbed in tanks contain- ing hydrochloric acid or sulphuric acid. This solution upon the addi- tion of an alkali gives up its ammonia, which is dissolved in distilled water, forming thereby the ammonium hydroxide or aqua ammonia of commerce. Ammonia is sometimes prepared from the residues of the beet sugar industry, from the refuse of slaughter houses and tanneries, and from the gases from coke ovens. It is not obtained directly from the nitrogen of the air. Properties of Ammonia. Ammonia gas is colorless. It has an exceedingly pungent odor, and if inhaled sud- denly or in large quantities it brings tears to the eyes and may cause suffocation. It is a light, volatile gas, being only .59 times as heavy as air. A liter of the gas at o and 760 mm. weighs .77 gm. It will not burn in the air, nor will it support the combustion of a blazing stick ; but if the air is heated or if its proportion of oxygen is increased, a jet of ammonia gas will burn in it with a yellowish flame, thereby illustrating the. broader application of the term combustion. Ammonia gas is easily liquefied if reduced to oC. and subjected to a pressure of 4^ atmospheres, while at 34 C. it liquefies at the ordinary atmospheric pressure. Compounds of Nitrogen. 149 Liquefied ammonia is often called anhydrous ammonia, because it contains no water. It boils at 33. 5 C. Hence, if it is exposed to the air or warmed in any way, it changes back to a gas, and in so doing absorbs considerable heat. This fact has led to the extensive use of liquid ammonia in the manufacture of ice. Ammonia is a strong alkali, and was called formerly the volatile alkali. Priestley, who discovered and studied the gas, called it alkaline air. Another marked property of ammonia gas is its solu- bility in water. A liter of water at oC. dissolves 1148!. of gas (measured at OC. and 760 mm.), and at the ordinary temperature I 1. of water dissolves about 700 1. of gas. This solution of the gas is usually called ammonia, though other names, especially ammonium hydroxide, are sometimes applied to it. Commercially it is known as aqua ammonia, ammonia, or ammonia water. It gives off the gas freely, when heated, as may easily be discovered by the odor or by the formation of the dense white fumes of ammonium chloride (NH 4 C1) when the solution is ex- posed to hydrochloric acid. The solution is lighter than water, its specific gravity being about .88, and contains about 35 per cent (by weight) of the gas. It is a strong alkali a caustic alkali, neutralizes acids and forms salts, and acts in many respects like sodium hydroxide. Ammonium Hydroxide and Ammonium Compounds. When ammonia gas is passed into water^it is believed that the ammonia combines with the water and forms a solution of an unstable compound having the formula NH 4 OH. This compound is ammonium hydroxide (or ammonium hydrate). Its formation may be represented thus NH 3 -f- H 2 O NH 4 OH Ammonia Water Ammonium Hydroxide 150 Descriptive Chemistry. Ammonium -hydroxide acts like a base. It has a marked alkaline reaction ; it neutralizes acids and forms salts, thus NH 4 OH + HC1 NH 4 C1 + H 2 O Ammonium Chloride 2NH 4 OH + H 2 S0 4 = (NH 4 ) 2 S0 4 + 2 H 2 O Ammonium Sulphate These salts, ammonium chloride and ammonium sul- phate, have definite properties, and are strictly analogous to sodium salts. Thus, we have Sodium Salts Ammonium Salts Nad NH 4 C1 NaNO 3 * NH 4 NO 3 Na 2 S0 4 (NH 4 ) 2 S0 4 etc. etc. Hence, it is believed that ammonium compounds contain a group of atoms which acts like an atom of a metal. This group of atoms is called ammonium, and its formula is NH 4 . Ammonium has never been separated from its compounds, or if it has it is so unstable that it immedi- ately decomposes into ammonia gas and hydrogen. So also ammonium hydroxide has never been obtained free, for it decomposes readily into ammonia gas and water, thus NH 4 OH NH 3 + H 2 O Ammonium Hydroxide Ammonia Gas Water Ammonium is sometimes called a radical, because it is the root or foundation of a series of compounds. It is likewise called a hypotheti- cal metal, because its existence is assumed and it acts chemically like metals. Compounds of Nitrogen. 151 Ammonium Chloride is prepared by passing ammonia gas into dilute hydrochloric acid, by mixing ammonium hydroxide and hydrochloric acid, or by letting the two gases mingle. The equation for the essential reaction is NH 3 + HC1 = NH 4 C1 Ammonia Hydrochloric Acid Ammonium Chloride It is convenient to regard this compound as the ammonium salt of hydrochloric acid, as if it were formed by replacing the hydrogen of the acid by ammonium, just as sodium forms sodium chloride. Ammonium chloride is a white, granular or crystalline solid, with a sharp, salty taste. It dissolves easily in water, and in so doing lowers the temperature markedly. When heated to a high temperature it gradually breaks up into ammonia and hydrochloric acid. This kind of decomposition is called dissociation. Large quantities of ammonium chloride are made at one stage of the manufacture of ammonium hydroxide by passing the gas into hydro- chloric acid. The crude product is called " muriate of ammonia " to indicate its relation to muriatic (or hydrochloric) acid. It is largely used for charging Leclanche' batteries, as an ingredient of soldering fluids, in galvanizing iron, and in textile industries. The crude salt is purified by heating it gently in a large iron or earthenware pot, with a dome-shaped cover ; the ammonium chloride volatilizes easily and then crystallizes in the pure state as a fibrous mass on the inside of the cover, but the impurities remain behind in the vessel. The process of vapor- izing a solid substance and then condensing the vapor directly into the solid state is called sublimation. It differs from distillation in that the substance does not pass through an intermediate liquid state. The product of sublimation is called a sublimate. Sublimed ammonium chloride is known as sal ammoniac. Ammonium Sulphate is made by passing ammonia gas into sul- phuric acid, or by adding ammonium hydroxide to the acid, thus 2NH 4 OH + H 2 SO 4 = (NH 4 ) 2 SO 4 + 2 H 2 O Ammonium Hydroxide Ammonium Sulphate 152 Descriptive Chemistry. * The commercial salt is a grayish or yellowish solid. It is used as a Constituent of fertilizers, since it is rich in nitrogen, and in making ammonium alum and other ammonium compounds. Ammonium Nitrate is made by passing ammonia into nitric acid, or by allowing ammonia gas and the vapor of nitric acid to mingle, thus NH 3 + HNO 3 = NH 4 NO 3 Ammonia Nitric Acid Ammonium Nitrate It is a white salt which forms beautiful crystals. It dissolves easily in water with. a fall of temperature. Its chief use is in the preparation of nitrous oxide (see this compound). Ammonium Carbonate is an impure salt as found in commerce, being a mixture of acid ammonium carbonate (HNH 4 CO 3 ) and a related compound. When pure and fresh it is transparent, but on ex- posure to the air it loses ammonia and turns white. It is used to pre- pare some kinds of baking powder, to scour wool, as a medicine, and to prepare smelling salts, since it gives off ammonia readily. Other ammonium compounds are sodium ammonium phosphate or microcosmic salt (HNaNH 4 PO 4 ), ammonium sulphocyanate (NH 4 SCN), and ammonium sulphide ( (NH 4 ) 2 S). Uses of Ammonia. Ammonia in the different forms is widely used as a cleansing agent, especially for the re- moval of grease, as a restorative in cases of fainting or of inhaling irritating gases, in dyeing and calico printing, and in the manufacture of dyestuffs, sodium carbonate, and ice. Its salts have many domestic, industrial, and agri- cultural uses. The Use of Ammonia as a Refrigerant and in making Ice depends upon the fact that many liquids in passing into a gas absorb heat. Liquefied ammonia (not the ordi- nary liquid ammonia) changes rapidly into a gas when its temperature is raised or the pressure reduced. Hence, if anhydrous ammonia is allowed to flow through a pipe sur- rounded by brine, the ammonia evaporates in the pipe and cools the brine, which may be used as a refrigerant or for Compounds of Nitrogen. 153 making ice. In some cold storage houses, breweries, packing houses, and sugar refineries, this cold brine is pumped through pipes placed in the rooms where a low temperature is desired. The construction and operation of an ice-making plant are essentially as follows : Liquefied ammonia is forced from a tank into a series of pipes which are submerged in an immense vat filled with brine. Large galvanized iron cans containing pure water to be frozen are immersed in the brine, which is being kept below the freezing point of water by the rapid evap- oration of the ammonia in the pipes. In about sixty hours the water in the cans is changed into a cake of ice weighing about three hundred pounds. As fast as the ammonia gas forms in the pipes, it is removed by exhaust pumps into another tank, where it is recondensed to liquefied ammonia and conducted, as needed, into the first tank to be used again. The ammonia is thus used over and over without appreciable loss. The pure water is sometimes obtained by condensing the exhaust steam from the boilers used to operate the machinery, though it usually comes from a deep well. Most ocean steamers have an ice plant, and in large cities in warm climates manufactured ice is a com- mon commodity. Composition of Ammonia Gas. Numerous experiments show that ammonia gas has the composition expressed by the formula NH 3 . (1) Dry ammonia gas passed over heated magnesium decomposes into hydrogen and nitrogen. The hydrogen may be collected and tested, but the nitrogen combines with the magnesium, forming a yellowish green powder called magnesium nitride, thus 2NH 3 + 3 Mg = Mg,N 2 + 3 H 2 Magnesium Magnesium Nitride These facts show that ammonia contains nitrogen and hydrogen. (2) If a bottle is filled with chlorine gas and plunged mouth downward into a vessel containing ammonium hydroxide, dense white fumes fill the bottle, the greenish chlorine gas disappears, and the liquid rises in the bottle ; after the bottle has stood mouth downward in a dish con- taining dilute hydrochloric acid (to neutralize the excess of ammonia), the gas in the bottle will be found to be nitrogen. The chlorine with- 154 Descriptive Chemistry. draws the hydrogen from the ammonia of the ammonium hydroxide, leaving the nitrogen free, thus NH, + 3 Q - N + 3 HC1 Ammonia Chlorine Nitrogen Hydrochloric Acid (3) The same experiment, if performed accurately, shows that one volume of nitrogen combines with three volumes of hydrogen to form ammonia gas. A tube containing a known volume of chlorine is pro- vided with a funnel through which concentrated ammonium hydroxide is dropped into the chlorine, until the reaction ceases (Fig. 22). After the excess of ammonia is neutralized with sulphuric acid, the volume of nitrogen left is one third of the original volume of chlorine gas. Now hydrogen and chlorine com- bine in equal volumes, hence the volume of hydrogen withdrawn from the added ammonia must be equal to the original volume of chlorine. But this volume is three times the volume of nitrogen, therefore there must be three times as much hydrogen as nitrogen in ammonia gas. (4) When electric sparks are passed through ammonia gas, it is decomposed into nitrogen and hydrogen. Now if oxygen is added, and an electric spark passed through the mixture, the oxygen and hydrogen combine. The volume of the remaining nitrogen is one fourth of the mixture of nitrogen and hydrogen, hence the hydrogen must have been three fourths ; that is, the volume of FIG. 22. Appa- hydrogen in the original volume ammonia was three ratus for determin- t i mes t i iat o f the nitrogen. (5) The gravimetric iner the composition ... r . c , , .,. . jr composition or ammonia gas is found by oxidizing it, and weighing the water and nitrogen, which are the only products. The result shows that fourteen parts of nitrogen combine with three parts of hydrogen. (6) The vapor density has been found to be 8.5. These facts require NH 3 as the simplest formula for ammonia and 17 as its molecular weight. Independent experiments verify this molecular weight. NITRIC ACID. Nitric Acid is one of the most useful compounds of nitrogen. It was known to the alchemists, who used it Compounds of Nitrogen. 155 to prepare a mixture which dissolves gold. Nitric acid is used in the preparation of many nitrogen compounds. Formation of Nitric Acid. When moist animal or vegetable matter containing nitrogen decays in the presence of an alkali, nitric acid is formed ; it is neutralized at once by the alkali, so nitrates salts of nitric acid are the final products. This chemical change is known as nitri- fication, and it is caused, or largely influenced, by minute living organisms called bacteria. Nitrification is constantly going on in the soil and is an exceedingly helpful process, since it transforms harmful waste matter into valuable plant food. As a result of nitrification, there are vast deposits of nitrates, espe- cially in desert regions and tropical countries. For example, potassium nitrate (KNO 3 ) is found in the soils near large cities in India, Persia, and Egypt. Nitric acid is formed in small quantities when electric sparks are passed through moist air. Hence nitric acid or its salts can be detected in the atmosphere after a thunder- storm. This chemical change is now being applied on a large scale at Ni- agara Falls. Electric sparks are passed through confined air and the products are forced into a tower. Here they are absorbed in water or in a solution of lime ; thereby forming nitric acid or calcium nitrate. The latter is converted into sodium nitrate (see below). Preparation. Nitric acid is prepared in the laboratory by heating concentrated sulphuric acid with a nitrate, usu- ally sodium or potassium nitrate. About equal weights of nitrate and acid are put into a glass retort and gently heated. The nitric acid distils into a receiver, which is kept cool by running water, ice, or moist paper. The Descriptive Chemistry. chemical change at a low temperature is represented by the equation NaNO 3 + H 2 SO 4 = HNO 3 + HNaSO 4 Sodium Nitrate Sulphuric Acid Nitric Acid Acid Sodium Sulphate 85 +98 = 63 + 120 But if the temperature is high and an excess of the nitrate is present, the equation is 2NaNO< H 2 SO 4 = 2HNO 3 Na 2 SO 4 170 +98 126 4- 142 A high temperature, however, decomposes part of the nitric acid, hence excessive heat is usually avoided. FlG. 23. Apparatus for the manufacture of nitric acid. Nitric acid is manufactured on a large scale by heating sodium nitrate and sulphuric acid in a large cast-iron retort (A) connected with huge glass or earthenware bottles (Z?, B, ), arranged as shown in Figure 23 ; the last bottle is connected with a tower filled with coke over which water trickles to absorb the vapors which escape from the bottles. The acid vapors are also often absorbed in earthenware or glass tubes. Properties. Pure nitric acid is a colorless liquid, but the commercial acid is yellow or reddish, due to absorbed nitrogen compounds, chlorine, or iron compounds. It de- composes slowly in the sunlight or when heated, and a Compounds of Nitrogen. 157 brownish gas may often be seen in bottles of nitric acid. It absorbs water, and forms irritating fumes when exposed to the air. The specific gravity of the commercial acid is about 1.42, and it contains from 60 to 70 per cent of the real acid (HNO 3 ), the rest being water. If the water is removed by slowly distilling the commercial acid with concentrated sulphuric acid, the product contains from 94 to 99 per cent of the real acid and its specific gravity is about 1.51. When nitric acid is boiled, it loses either acid or water until the liquid contains approximately 68 per cent of nitric acid, and then it continues to boil unchanged at 120 C. Nitric acid is very corrosive. It turns the skin a perma- nent yellow color, and may cause serious burns. Many organic substances are turned yellow and sometimes com- pletely decomposed by it. It parts readily with its oxygen, especially when hot, and is therefore an energetic oxidizing agent. Charcoal burns brilliantly in hot acid, while straw, sawdust, hair, and similar substances are charred and even inflamed by it. Iron sulphide heated with nitric acid becomes iron sulphate, by the addition of oxygen, thus FeS + 2O 2 = FeSO 4 Iron Sulphide Oxygen Iron Sulphate Uses of Nitric Acid. Nitric acid is one of the com- mon laboratory acids. Large quantities are used in the manufacture of nitrates, dyestuffs, sulphuric acid, nitro- glycerine, gun cotton, in the refining of gold and silver, and in etching copper plates. Composition of Nitric Acid. Although the alchemists knew and valued nitric acid, its composition was a mystery until Lavoisier showed in 1776 that it contained oxygen and probably nitrogen. Its exact composition was determined by Cavendish in 1784-1785, by passing electric sparks through a mixture of oxygen and nitrogen in the pres- 158 Descriptive Chemistry. ence of water or caustic potash. The same facts had been observed, but not explained, by Priestley. Many independent experiments show that the composition of nitric acid is expressed by the formula HNO 3 . (1) When electric sparks are passed through a bottle containing moist air or a solution of potassium hydroxide, the water becomes acid to litmus or the liquid will be found to contain a trace of potassium nitrate. (2) Nitric acid may be reduced to ammonia by nascent hydrogen, thus showing that the acid contains nitrogen. (3) Conversely, if a mixture of ammonia and air is passed over a mass of hot, porous platinum, nitric acid is formed. (4) If the acid is allowed to flow through a hot porcelain or clay tube, oxygen is one of the gaseous products. Nitrates. Nitric acid is monobasic and forms a series of well-defined salts called nitrates. The interaction of nitric acid and most metals is exceedingly vigorous, and for this reason, probably, the alchemists called the acid aquafortis strong water. The reaction varies with the metal, strength of the acid, temperature, and the presence of resulting compounds. The solid product of the reaction is usually a nitrate, though some metals, such as tin and antimony, form oxides. The gaseous products are usually oxides of nitrogen, especially nitric oxide (NO), which, however, quickly forms nitrogen peroxide (NO 2 ) in the air. Hydrogen is never liberated so that it can be collected ; probably it immediately reduces the nitric acid to another compound of nitrogen. Nitrates are also formed by the action of nitric acid upon oxides, hydroxides, and carbonates, thus CuO Copper Oxide KOH Potassium Hydroxide Na 2 CO 3 Sodium Carbonate 2HNO = HN0 - Cu(NO,) 2 Copper Nitrate KN0 3 Potassium Nitrate 2HNO 3 = 2 NaNO 3 Sodium Nitrate CO H 2 H 2 O H 2 O Compounds of Nitrogen. 159 When nitric acid is poured upon copper, the liquid bub- bles violently and becomes hot, dense fumes of a reddish brown gas are given off, and the liquid turns blue owing to the dissolved copper nitrate. Other metals, such as zinc, iron, and silver, act in a similar way, though the nitrate is blue only in the case of copper. The usual equation for the chemical change with copper is 3Cu + 8HNO 3 = 3Cu(NO 3 ) 2 + 2 NO + 4 H 2 O Copper Nitrate Nitric Oxide When nitric oxide is exposed to the air, it changes at once into the reddish brown peroxide, thus NO + O NO 2 Nitric Oxide Oxygen Nitrogen Peroxide Nitrates as a rule are very soluble in water. They be- have in various ways when heated. Some, like sodium and potassium nitrates, lose oxygen and pass into nitrites ; others, like copper nitrate, form an oxide of the metal, an oxide of nitrogen, and oxygen ; and one, ammonium nitrate, decomposes into water and nitrous oxide (N 2 O). Since many nitrates, when heated, give up oxygen, they are powerful oxidizing agents. Potassium nitrate dropped on hot charcoal burns the charcoal vigorously and rapidly. This kind of chemical action is called deflagration. The Test for Nitrates (and of course for nitric acid) is as follows : Add to the solution of the nitrate a little concentrated sulphuric acid, and upon the cool mixture pour carefully a cold, dilute solution of fresh ferrous sulphate. A brown layer is formed where the two liquids meet. Nitrous Acid (HNO 2 ) has never been obtained in the free state, but its salts the nitrites are well known. Potassium nitrite (KNO 2 ) and sodium nitrite (NaNO 2 ) are formed by removing the oxygen from the corresponding nitrate by heating gently or by heating with lead. Nitrites give off brown fumes when treated with sulphuric acid, thus i6o Descriptive Chemistry. being readily di decomposition o amount in drin shed from nitrates. Nitrites are formed by the ' C matter, and the presence of a relatively large r indicates contamination by sewage. Aqua Regia is an old term which is still applied to a mixture of concentrated nitric and hydrochloric acids. The expression means "royal water," and indicates that the mixture dissolves gold and platinum the noble metals. Its solvent power depends mainly upon the free chlorine which is produced in the mixture by the oxidizing action of the nitric acid. The product of the action of aqua regia on metals is always the chloride of the metal. Oxides of Nitrogen. There are five oxides of ni- trogen : NAME. FORMULA. CHARACTERISTIC. Nitrous oxide N 2 O Colorless "as NO Colorless o'as Nitrogen trioxide Nitrogen peroxide NA NO 9 Blue liquid Brown gas Nitrogen pentoxide . . N,O, White solid j.^2^5 Only three of these are important, viz., nitrous and nitric oxides, and nitrogen peroxide. Nitrous Oxide is one of the numerous decomposition products of nitric acid, but it is usually prepared by decom- posing ammonium nitrate. This salt, if gently heated in a test tube provided with a delivery tube, first melts and then decomposes into water and nitrous oxide ; the gas may be collected over warm water. The equation of the chemical change is NH 4 NO 3 = N 2 O Ammonium Nitrate Nitrous Oxide 2H 2 O Compounds of Nitrogen. 161 This colorless gas has a sweet taste and a faint but pleas- ant odor. It is less soluble in hot than in cold water. The gas does not burn, but it supports the combustion of many burning substances, though not so vigorously as oxygen does. Sulphur, for example, will not burn in nitrous oxide, unless the sulphur is hot and well ignited at first. The most striking property of nitrous oxide is its effect on the human system. If breathed for a short time, it causes more or less nervous excitement, often manifested by laughter, and on this account the gas was called "laughing gas" by Davy. If breathed in large quantities, it slowly pro- duces unconsciousness and insensibility to pain. The gas is often used when insensibility is desired for a short time, as in dentistry. It is easily liquefied by cold and pressure, and is often used in this form to furnish the gas itself and to produce very low temperatures. It is a commercial article and is sold in small iron cylinders. Nitrous oxide was discovered by Priestley in 1776; but its composi- tion was not explained until 1799, when Davy, by an extensive study of its properties, proved it to be an oxide of nitrogen. In his enthusiasm Davy wrote a friend: "This gas raised my pulse upward of twenty strokes, made me dance about the laboratory as a madman, and has kept my spirits in a glow ever since." It is needless to say that the usual results are more quieting. The Composition of Nitrous Oxide is shown as follows : By ex- oloding equal volumes of nitrous oxide and hydrogen, only nitrogen Remains, and its volume equals the original volume of nitrous oxide. The oxygen unites with the hydrogen to form water, and there is just enough oxygen to unite with a volume of hydrogen equal to the volume of the nitrous oxide. Therefore, the oxygen in the nitrous oxide must have been equal to half the volume of the nitrogen, since oxygen and hydrogen combine in the ratio of one to two. Furthermore, experiment has shown that the weights of equal volumes of nitrous oxide and ni- trogen are in the ratio of 44 to 28. Therefore, the smallest part of oxygen united with the nitrogen must weigh 16 ; and since the nitrogen weighs 28, the formula must be N.,0. 1 62 Descriptive Chemistry. Nitric Oxide has long been known, since it is the usual gaseous product of the interaction of nitric acid and metals. It is usually prepared by the interaction of copper and dilute nitric acid (sp. gr. 1.2). The equation for the com- plex chemical change is usually written thus 3Cu + 8HNO 3 = 2 NO + Cu(NO 3 ) 2 + 4 H 2 O Copper Nitric Acid Nitric Oxide Copper Nitrate The gas thus prepared is impure, and it is customary to use ferrous sulphate and nitric acid as a source of the pure gas. Nitric oxide is a colorless gas, but upon exposure to the air, it combines at once with oxygen, forming dense red- dish brown fumes of hydrogen peroxide. The simplest equation for this change is NO + O = NO 2 Nitric Oxide Nitrogen Peroxide This property distinguishes nitric oxide from all other gases. It does not burn, nor does it support combustion unless the burning substance (e.g. phosphorus or sodium) introduced is hot enough to decompose the gas into nitro- gen and oxygen, and then, of course, the liberated oxygen assists the combustion. The Composition of Nitric Oxide is determined by heating iron or another metal in it. The oxygen of the oxide combines with the iron, and the nitrogen is left free. The resulting volume of nitrogen is half the volume of the nitric oxide taken. Hence nitric oxide contains equal volumes of nitrogen and oxygen. By an independent experiment the molecular weight is found to be 30. Hence the formula must be NO. Nitrogen Peroxide is the reddish brown gas formed by the direct combination of nitric oxide and oxygen. Thus NO + O NO 2 Nitric Oxide Nitrogen Peroxide Compounds of Nitrogen. 163 It is also produced by heating certain nitrates. Thus Pb(NO 3 ) 2 = 2NO 2 + PbO + O Lead Nitrate Nitric Oxide Lead Oxide Oxygen The fumes of nitrogen peroxide always appear when nitric acid and metals interact, but, as already stated, the fumes are not produced at first, being the result of a second chemical change when the real product, nitric oxide, comes in contact with oxygen of the air. Nitrogen peroxide is poisonous. It dissolves in water ; it also dissolves in concentrated nitric acid, forming fuming nitric acid. At very low temperatures nitrogen peroxide is a colorless solid. At about 10 C. it is a yellowish liquid, and as the temperature rises the color grows darker, until at 22 C. the liquid boils and gives off the familiar reddish brown gas. Above 140 C. this gas begins to lose its color, and at 600 C. the color entirely disappears. The density of the gas at low temperatures indicates the formula N 2 O 4 , whence the name nitrogen tetroxide, often used. But the density at about 140 C. indicates the formula NO 2 . Nitrogen Trioxide, N 2 O 3 , and Nitrogen Pentoxide, N 2 O 5 , are unstable compounds and have no practical importance. They are the anhy- drides of nitrous and nitric acids, thus N 2 8 + H 2 O 2HNO 2 Nitrogen Trioxide Nitrous Acid N 2 5 + H 2 O 2HNO ? Nitrogen Pentoxide Nitric Acid EXERCISES. 1. Name several sources of ammonia gas. How is ammonia gas prepared in the ' laboratory ? Give the equation for the reaction. State its important properties. 2. What is ammonium hydroxide ? How is it prepared on a large scale ? Summarize its properties. What are its uses ? 3. What is the meaning and significance of (a) volatile alkali, 164 Descriptive Chemistry. ($) anhydrous ammonia, (V) spirits of hartshorn, (d) sal volatile, (^) muriate of ammonia, (/") sal ammoniac, (g) aqua for'tis ? 4. Why is NH 3 the formula of ammonia gas ? 5. Give several tests for (a) ammonia, and (V) nitric acid. 6. What different meanings may the word ammonia have ? What is ammoniacal liquor? Gas liquor? Aqua ammonia? Ammonium hydrate ? Ammonia of commerce ? Ammonia water ? 7. How is ammonia gas liquefied ? Describe the manufacture of ice by liquid ammonia. 8. Develop the topics : (a) ammonium is a radical ; () nitric acid is an oxidizing agent ; (<:) nitrates are unstable ; (d) fuming nitric acid. 9. Give the formula, method of preparation, properties, and uses of (a) ammonium chloride, (b) ammonium nitrate, (c} ammonium sulphate, (d) ammonium carbonate. 10. How is nitric acid formed (a) in the soil, () in the air ? How is it prepared (a) in the laboratory, (b) on a large scale ? Summarize (a) the physical properties of nitric acid, and (b} its chemical properties. For what is it used ? 1 1 . What is the formula of nitric acid ? Summarize the evidence of its composition. 12. What are nitrates ? How are they formed ? What is the effect of heat upon () potassium nitrate, ($) copper nitrate, (c} am- monium nitrate ? Give other properties of nitrates. What is the test for nitrates ? 13. What are nitrites ? How are they formed ? How are they distinguished from nitrates ? 14. What is aqua regia? For what is it used ? Why so called ? What is the chemical action of aqua regia on gold ? Upon what prop- erty of nitric acid does its chemical action depend ? 15. Give the names and formulas of the five oxides of nitrogen. Describe the preparation of nitrous oxide. State briefly its properties. For what is it used ? Who discovered it ? What did Davy call it ? Why ? Summarize the evidence of the composition of nitrous oxide. 1 6. Describe the preparation of nitric oxide. State the equation for the reaction. What are its properties ? 17. How is nitrogen peroxide prepared ? State its properties. How is it readily distinguished from all other oxides of nitrogen ? What two formulas have been given to nitrogen peroxide ? Why ? 1 8. What is (#) nitric oxide, (b) nitrous oxide, (c) nitrogen per- Compounds of Nitrogen. 165 oxide, (//) nitrogen tetroxide, (c) nitrogen trioxide, (d) nitrogen monoxide, (e) nitrogen pentoxide ? 19. State the equation for the preparation of (a) nitric acid at a low temperature, () nitric acid at a high temperature, (V) ammonium chloride, (d) ammonium hydroxide from water and ammonia, (d} ni- trous oxide, (e) nitrogen peroxide, (/") copper nitrate. 20. Define and illustrate () sublimation, ($) sublimate, (V) nitrifi- cation, (d} deflagration, (>) nitrate, (/") ammonium compound. 21. What is the valence of nitrogen in ammonia gas ? In ammo- nium ? In ammonium hydroxide ? 22. (a) Why are there no acid nitrates ? () What is the valence of nitrogen in nitric acid, copper nitrate, nitrous oxide, nitric oxide, nitrogen peroxide, nitrogen trioxide, nitrogen pentoxide ? PROBLEMS. 1. How many grams of ammonia gas can be obtained from 2140 gm. of ammonium chloride by heating with lime ? 2. Calculate the percentage composition of (a) ammonium chloride, () ammonium hydroxide, (Y) ammonium sulphate, (d} ammonium nitrate. 3. Calculate the simplest formula of the compounds having the per- centage composition (a) N = 82.35, H = 17.64; and () N = 26.17, Cl = 66.35, H = 7.48. 4. Calculate the percentage composition of (a) nitric acid, () po- tassium nitrate (KNO 3 ), (V) sodium nitrate. 5 . How many grams of nitric acid can be obtained by heating a kilogram of sodium nitrate with sulphuric acid at a low temperature ? 6. If the specific gravity of a sample of nitric acid is 1.522, (a) what will 100 cc. weigh, and (b) what volume must be taken to weigh 100 grams ? 7. Calculate the simplest formula of the substances having the composition (a) O = 76.19, H = 1.58, N = 22.22; () N =13.86, K = 38.61, O = 47.52. CHAPTER XIII. PROPERTIES OF GASES GAY-LUSSAC'S LAW OF GAS VOLUMES AVOGADRO'S HYPOTHESIS VAPOR DEN- SITY AND MOLECULAR WEIGHT MOLECULAR WEIGHTS AND ATOMIC WEIGHTS MOLECULAR FORMULA MO- LECULAR EQUATIONS VALENCE. Properties of Gases. Extensive study of gases shows that they all conform to simple laws. Thus we have already seen that they behave uniformly with changes of pressure (Boyle's law) and with changes of temperature (Charles's law). Other simple relations prevail. Gay-Lussac's Law. Gases combine by volume in simple ratios. Experiment has revealed the following facts about the COMBINATION OF GASES BY VOLUME. VOLUMES OF COMPONENTS. VOLUMES OF PRODUCTS. 2 vol. hydrogen I vol. oxygen 2 vol. water vapor I vol. chlorine i vol. hydrogen 2 vol. hydrochloric acid gas 3 vol. hydrogen i vol. nitrogen 2 vol. ammonia gas 2 vol. nitrogen i vol. oxygen 2 vol. nitrous oxide gas 2 vol. nitrogen 3 vol. oxygen 2 vol. nitrogen trioxide gas 1 66 Avogadro's Hypothesis. 167 Additional illustrations will be given in later chapters. The simple ratio which exists between the gas volumes, whether components or products, has been found to be true of all gases. The law was pointed out in 1808 by Gay-Lussac, who stated the relation substantially as follows: Gases combine in volumes which bear a simple ratio to each other and to that of the product. By " a simple ratio " we mean one made up of small whole numbers. As a rule, the product occupies two unit volumes. "" Avogadro's Hypothesis. In 1811 an Italian physi- cist proposed an hypothesis to account for the similar behavior of gases. At that time the properties of gases were not generally known, and the views of Avogadro were overlooked until about 1860. Since then the hypo- thesis has been helpful in explaining many facts, and it is generally accepted by chemists as a very probable assumption. It may be stated thus: There is an equal number of molecules in equal vohimes of all gases at the same temperature and pressure. This statement cannot be proved directly by experiment, but there is much physical, chemical, and mathematical evidence in harmony with it. According to Avogadro's hypothesis a liter of hydrogen and a liter of oxygen at the same temperature and pres- sure contain the same number of molecules, though we do not know how many. Suppose, however, that each liter contained 1000 molecules. A liter of hydrogen weighs 0.0896 gm. and a liter of oxygen at the same tem- perature and pressure weighs 1.43 gm. But 0.0896 and 1.43 are in the same ratio as i and 16. Therefore, since 1 68 Descriptive Chemistry. a thousand molecules of oxygen weighs 16 times more than a thousand molecules of hydrogen, a single molecule of oxygen must weigh 16 times more than a single molecule of hydrogen. Therefore, in general, in order to find how much heavier any gaseous molecule is than a hydrogen molecule, it is only necessary to compare the weights of equal volumes of hydrogen and the gas under examination. An application of Avogadro's hypothesis is made in course of the following argument, which proves that a molecule of hydrogen consists of two atoms : One volume of hydrogen combines with one volume of chlorine to form two volumes of hydrochloric acid gas. Suppose the volume of hydrogen contained 100 molecules. Then, according to Avogadro's hypothesis, the equal volume of chlorine will contain 100 molecules, while the two volumes of the product will contain 200 molecules of hydrochloric acid gas. That is 100 molecules of Hydrogen + 100 molecules of Chlorine = 200 molecules of Hydrochloric Acid Gas. Now every molecule of hydrochloric acid gas contains at least one atom each of hydrogen and chlorine, and the 200 molecules must contain 200 atoms each of chlorine and hydrogen. Therefore each molecule of hydrogen and of chlorine must be divisible into two atoms, since the 100 hydrogen and the 100 chlorine molecules provide the 200 hydrogen atoms and the 200 chlorine atoms in the 200 molecules of hydrochloric acid gas. Similar reasoning leads to the conclusion that the molecules of oxygen, nitrogen, and most elementary gases consist of two atoms. Vapor Density and Molecular Weight. It was stated in a previous chapter that a molecular weight is the sum of the weights of the atoms in the molecule. But this method of finding the molecular weight is useless, unless we first know the formula, and in many cases the formula cannot be chosen until after the molecular weigjit has been found by several methods. Hence, the determination of molecular weights is an important matter. In the case Vapor Density and Molecular Weight. 169 of gaseous or volatile elements and compounds, it is often accomplished by finding the vapor density of the substance. There is a direct and simple relation between molecular weight and vapor density. By vapor density we mean the ratio of the weight of a gas to the weight of an equal volume of hydrogen at the same temperature and pressure. Thus, the vapor density of steam is 9, because experiment shows that it weighs 9 times more than an equal volume of hydrogen under the same conditions of temperature and pressure. Therefore the molecular weight of steam is 9 times the molecular" weight of hydrogen. But the molec- ular weight of hydrogen is 2, since its molecule contains two atoms each weighing I. Therefore, the molecular weight of steam is 18, or twice the vapor density. The general fact that the molecular weight of a gaseous com- pound is twice its vapor density is clearly seen from the following table showing the RELATION BETWEEN VAPOR DENSITY AND MOLECULAR WEIGHT. GAS. VAPOR DENSITY. MOLECULAR WEIGHT. Carbon dioxide . Ammonia. .... 22 8.5 44 17 Hydrochloric acid Water vapor (steam) . 18.25 9 36.5 18 Hence, a determination of the vapor density of a com- pound or an element allows us to select the correct molec- ular weight and assign the proper formula. The vapor densities of the elements mercury and cadmium show that the atom and molecule are identical, while the vapor densities of phosphorus and arsenic indicate that the molecule of each consists of four atoms. A molecule of oxygen contains two atoms, but a molecule of ozone contains three ; therefore, the formula of ozone is O 3 . i jo Descriptive Chemistry. Other Methods of determining Molecular Weights. Some sub- stances cannot be vaporized without decomposition. The molecular weights of such substances cannot, of course, be found by the vapor density method. If a substance dissolves without decomposition, its molecular weight can be determined by the boiling-point or freezing- point method, which was briefly described in Chapter X. The above methods give approximate results. Exact molecular weights are found by accurate quantitative analysis. Suppose we wished to find the molec- ular weight of acetic acid. Silver acetate is analyzed and found to con- tain 64.65 per cent of silver ; the per cent of the remaining elements of the molecule must be 35 35. The atomic weight of silver is 107.93, if the atomic weight of oxygen is 16. Hence, the weight of the silver acetate molecule, except the silver, is found by the proportion 107.93: x : : 64.65 : 35.35. x= 59.02. Silver acetate is formed by replacing one atom of the hydrogen of the acid by one atom of silver. Therefore, the weight of the molecule of acetic acid is found by adding to 59.02 the weight of one atom of hydro- gen. That is, the exact molecular weight of acetic acid is 60.028 (i.e. 59.02 + 1.008). Determination of Atomic Weights. The atomic weight of an element, as already stated, is a relative weight. It is a number expressing the relation of the weight of an atom of a given element to the weight of an atom of some element chosen as a standard. Thus, if we say that the atomic weight of nitrogen is 14, we mean that the relation between the weight of the nitrogen atom and that of the hydrogen atom is 14 to i, if we adopt the hydrogen atom as the standard atom ; or we mean that the relation between the weight of the nitrogen atom and that of the oxygen atom is 14 to 16, if we adopt the oxygen atom as the standard. The approximate atomic weights are usually expressed in round numbers, and do not vary much with the standard. Wherever exact atomic weights are used in this book, the oxygen standard is the basis. Determination of Atomic Weights. 171 In Chapter IX it was stated that the determination and selection of atomic weights are based on several principles. This subject can now be appropriately considered. One method of selecting the atomic weight is illustrated by the case of chlorine, which has the atomic weight 35.5. The molecular weights of several chlorine compounds are found by the vapor density method. The compounds are analyzed to find the number of grams of chlorine in the number of grams of the compound equal to the determined molecular weight. And the highest common factor of these weights of chlorine is taken as the atomic weight of the element. A concise view of the method is shown in the following TABLE OF CHLORINE COMPOUNDS. COMPOUND. MOLECULAR WEIGHT. WEIGHT OF CHLORINE. H. C. F. Hydrochloric acid .... Chlorine peroxide .... Cyanogen chloride .... Chlorine sas 36.5 67.5 6l. S 71 35-5 35-5 35-5 71 i x 35-5 i x 35-5 i x 35.5 2 X 3C C Chlorine monoxide . . . Phosphorus trichloride . . Chloroform 8? 137.5 I IQ.1 71 106.5 io6.cr * JJ'J 2 X 35.5 3 x 35.5 ? x "K.c( Carbon tetrachloride . . . 170 142 4 x 35.5 Thirty-five and five tenths is therefore selected as the approximate atomic weight of chlorine. Atomic weights can also be determined by analysis if we know the proportion in which the atoms combine to form a molecule of the compound analyzed. Thus, the Belgian chemist, Stas, who made masterly determinations of atomic weights, found that 121.4993 gm. of silver 172 Descriptive Chemistry. chloride were formed by burning 91.462 gm. of silver in chlorine. He knew that one atom of silver and one of chlorine unite to form silver chloride ; he also accepted 35.453 as the atomic weight of chlorine. Hence, he calcu- lated the atomic weight of silver thus 121.4993-91.462= 30.0373, which is the weight of the chlorine used. Therefore 91.462 : 30.0373 : : x : 35-453, * = 107.95, the atomic weight of silver. Approximate atomic weights of the solid elements, espe- cially the metals, are checked by applying the law of spe- cific heats. This law was announced by Dulong and Petit in 1819. It is stated as follows :- The product of the specific heat and atomic weigJit of tJie solid elements is a constant quantity. By Specific heat we mean the quantity of heat necessary to raise the temperature of a substance one degree com- pared with the quantity necessary to raise the temperature of the same weight of water one degree. If the same quantity of heat is imparted to equal weights of water and mercury, the temperature of the mercury will be much higher about 32 times higher than that of the water. That is, the mercury requires only about ^ as much heat as the water. In other words, the specific heat of mercury is ^2> or o -3 r 2 - The specific heat of other elements is simi- larly found. The constant quantity found by multiplying the specific heat by atomic weight is approximately 6.25. This rela- tion is illustrated by the following Determination of Atomic Weights. 173 TABLE OF SPECIFIC HEATS. ELEMENT. SPECIFIC HEAT. ATOMIC WEIGHT. PRODUCT. Calcium O I7O AQ 6 8 Copper O OQC 6^ 6 6 04. Iron '^yj O.I 14. c6 638 Lead Potassium . Sodium .... 0.031 0.166 O.2Q'? 207 39 2"J 6.41 6.47 6 7^ Sulphur 0.178 M 57 Tin O OCC I IQ / 6 cj. Zinc ->.W-J} O OQ4. 6c A 6 i c ' w y i f ^y't The use of this law in checking atomic weights may be illustrated as follows : The specific heat of silver is found by experiment to be 0.057; if 6.25 is divided by this num- ber, the quotient is approximately 109. This result agrees approximately with 108 the accepted atomic weight of silver. Again, the specific heat of mercury is 0.0312; if 6.25 is divided by this number,- the quotient, 200, indicates that the atomic weight of mercury is 200 a value obtained by other methods. This law has been of assistance in the final selection of the approximate atomic weight of several elements. Thus, the atomic weight of uranium was finally accepted as about 238 instead of 119. Both values agreed with analyses, but only the former conformed to Dulong and Petit's law. The plan followed in determining the atomic weight of zinc illustrates the methods actually used. (a) When zinc interacts with dilute hydrochloric or sulphuric acid, hydrogen is liberated ; and if a known weight of zinc is used, the weight of zinc needed to liberate I gm. of hydrogen is easily calculated. This number, as we have already seen, is the equivalent of zinc (see Equivalents, Chapter IX). Now if one atom of zinc replaces one atom 174 Descriptive Chemistry. of hydrogen, then the atomic weight of zinc and the atomic weight of hydrogen will have the same ratio as the weight of zinc and the weight of hydrogen found by experiment. According to experiment the equivalent of zinc is about 32.5. This is its relation, atom for atom, to hydrogen, and, thus far, is its atomic weight. () When zinc and hydrochloric acid interact, zinc chloride is formed. If it is analyzed, the proportion of zinc to chlorine is about 32.5 to 35.5. If the elements combine, atom for atom, the atomic weight of zinc is 32.5 (assuming that 35.5 is the atomic weight of chlorine). (<:) When zinc is burned in air, zinc oxide is formed. If this com- pound is analyzed, the proportion of zinc to oxygen is about 65 to 16. If the elements combine atom for atom, the atomic weight of zinc is about 65 (assuming that 16 is the atomic weight of oxygen). (df) According to these three determinations, the atomic weight of zinc is 32.5 or 65. We have assumed that the elements unite atom for atom in each compound. This is an incorrect assumption, because an atom of zinc cannot have two different weights 32.5 and 65. If the atomic weight is 32.5, zinc oxide must consist of one atom of oxygen and two of zinc. But if the atomic weight is 65, zinc chloride must consist of two atoms of chlorine and one of zinc, and two atoms of hydrogen must have been replaced by one of zinc. (e) The molecular weight of zinc chloride is found by the vapor density method to be about 133. If zinc chloride consists of two atoms of chlorine and one of zinc (weighing 65), its molecular weight is about 136. In other words, it is evident that our assumption regard- ing the number of atoms in zinc chloride is highly probable. CO We are not absolutely positive, however, that the zinc in a molecule of zinc chloride may not be one atom weighing 65, or two atoms weighing 32.5 each. But the atomic weight of zinc determined by applying the law of specific heats is 664 (i.e. 6.25 -:- 0.094). This shows clearly that the atomic weight of zinc is approximately 65. Molecular Formula. In Chapter IX a method was given for finding the simplest formula of a compound, viz., by dividing the percentage of each element by its atomic weight. But the simplest formula is not always the mo- lecular formula ; that is, it does not always express the composition and number of atoms in a molecule of the Molecular Equations. 175 compound in the gaseous state. Every formula, however, is designed to be a molecular formula. Since the molecu- lar weight of a compound is twice its vapor density, the molecular formula can be calculated from the simplest formula. Thus, the simplest formula of a compound of carbon and hydrogen was found to be CH 2 . Its vapor density was found to be 81.4. Hence its molecular weight must be 162.8, which is nearly twelve times that corre- sponding to CH 2 . Therefore the molecular formula is C 12 H 24 . Molecular formulas of other compounds may be similarly found. Molecular Equations. Equations which represent re- actions between gases are sometimes written as molecular equations. Such equations represent changes as taking place between the smallest possible physical units, that is, between molecules. The molecular equation for the for- mation of water from hydrogen and oxygen is 2 H 2 + O 2 = 2 H 2 O. It is read thus : Two molecules of hydrogen unite with one molecule of oxygen to form two molecules of water. Since most elementary gases consist of molecules, such an equa- tion is strictly correct. It should be noted, however, that the proportions are the same as in the simpler form of the equation. For practical purposes the molecular equation is preferable only in the case of gases. Molecular equations are sometimes called volume or gas equations, because such equations tell at a glance the volumes involved in the re- action. Thus- H2 + Cl2 = 2HC1 means that one volume each of hydrogen and chlorine unite to form two volumes of hydrochloric acid gas. This equation is sometimes writ- ten H 2 + C1 2 = 2HC1 I VOL I VOl. 2 VOl. 176 Descriptive Chemistry. Valence. An examination of many formulas obtained by the principles just discussed shows certain regularities. Take, for example, some binary compounds of hydrogen. They fall into four groups, thus I. II. III. IV. HC1 H 2 O H 3 N H 4 C HBr H 2 S H 3 P H 4 Si Obviously, the atoms of these elements differ in their power of combining with hydrogen atoms. Some unite with one atom, some with two atoms, and so on. Atoms of other elements besides those in the above list differ in their combining power. The power of atoms of an ele- ment to hold in combination a certain number of other atoms is called the valence or quantivalence of the element. The valence of hydrogen is always one. Ele- ments which combine atom for atom with one atom of hydrogen have the valence one, and are called univalent elements or monads ; sodium and potassium are always univalent, and so is chlorine in hydrochloric acid. Ele- me'nts which combine with two atoms of hydrogen have the valence two, and are called bivalent elements or dyads ; oxygen, magnesium, and sulphur are bivalent elements. So, also, some elements like aluminium, are tri- valent or triads ; others, like carbon and silicon, are quadrivalent or tetrads; and some, like the nitrogen in nitric acid, are quinquivalent or pentads. Elements of the same valence combine with or replace each other atom for atom. Thus, one atom of sodium replaces one atom of hydrogen in hydrochloric acid ; and one atom of oxygen combines with one atom of magnesium. Elements of dif- ferent valence form compounds in which, as a rule, the number of atoms is such that the valences balance, Thus, Valence. 177 a dyad combines with two monads (as in H 2 O), a triad with three monads (as in NH 3 ), two triads with three dyads (as in A1 2 O 3 ), one tetrad with two dyads (as in CS 2 ), and so on. Such compounds, in which the capacity for further union has ceased, are said to be saturated or to have no free bonds. Compounds in which the valence is not bal- anced, or in which free bonds exist, are called unsaturated (see Ethylene). The valence of an element is always the same in the same compound, but it often varies. Thus, the valence of ni- trogen is one in N 2 O, two in NO, three in N 2 O 3 , four in NO 2 , and five in HNO 3 . Hydrogen, as stated above, always has a valence of one ; it is also believed that the valence of oxygen is always two. If an element forms no hydrogen compound, its valence is determined from compounds con- taining elements which are univalent, such as chlorine, bromine, and sodium. The valence of elements in saturated compounds of two elements is easily deduced from the formula, because in such compounds the total valence of all the atoms of each element must be equivalent Thus in the formula CaO, the valence of calcium is two, because the single atom of calcium is combined with a single atom of a bivalent element. The valence of phosphorus in P 2 O 5 is five, be- cause the two atoms furnish a total valence of ten, which is required by the five atoms of the bivalent element oxygen. In CH 4 the valence of carbon is four, because the single atom is combined with four atoms of hydrogen. Radicals have a valence, since in chemical changes they act like atoms. The valence of ammonium (NH 4 ) is one, and of hydroxyl (OH ) is one. Thus, NH 4 C1 is the formula of ammonium chloride, NaOH of sodium hydroxide, but Ca(OH) 3 of calcium hydroxide. ' 1 7 8 Descriptive Chemistry. The valence of elements in unsaturated compounds can- not be told by mere inspection ; a knowledge of the prop- erties of the compound is necessary. So also the valence of some elements in compounds containing three or more elements is not readily told from the formulas^ some knowledge of the methods of formation, relations to other compounds, and general properties is needed. A discus- sion of these principles is beyond the scope of this book. However, in the case of most acids, bases, and salts, an arbitrary rule may be cited. In these compounds the total valence of the oxygen atoms balances the total valence of the other elements. Thus, in nitric acid, HNO 3 , the va- lence of nitrogen is nve, while in nitrous acid, HNO 2 , it is three. Some chemists prefer to regard valence as the quotient obtained by dividing the atomic weight by the equivalent weight. For example, the valence of oxygen is 2 the quotient of 16 -4- 8. Such a view is not inconsistent with the one generally held, because valence is the direct outcome of. composition. The valence of elements may be represented in several ways, e.g. H', H , O , O = , N . Sometimes formulas are written to show the valence, e.g. / H Hydrochloric acid, H - Cl, Water, H - O - H, Ammonia, N - H. Such formulas are called structural or graphic formulas to distinguish them from the ordinary or empirical formulas. Structural formulas are not intended to show how the atoms are arranged in space. We know very little about the space relations of atoms. They simply indi- cate certain relations not shown by the empirical formulas. They are especially helpful in organic chemistry (see Chapter XXXI). EXERCISES. i . Review (a) Boyle's law, and (ft) Charles's law. 2. State and illustrate Gay-Lussac's law. 3. Give a brief account of (a) Gay-Lussac ? () Avogadro, (c) Stas, Exercises. 179 4. State and illustrate Avogadro's hypothesis. 5. What is the relation of the molecular weight of a gas to (a) the molecular and () the atomic weight of hydrogen ? 6. (a) State the argument proving that a molecule of hydrogen con- sists of two atoms, (b} Apply the same argument to oxygen. 7. What is the relation between molecular weight and vapor den- sity ? Illustrate your answer. What application is made of this relation ? 8. Why is the formula of water H 2 O and not HO or H 2 O 2 ? 9. Why is the formula of ozone O 3 ? 10. (a) How are molecular weights determined ? () How are atomic weights found from molecular weights ? 11. Illustrate the method of determining atomic weights by chemical analysis. 12. What is a molecular formula ? What is the molecular formula of oxygen, nitrogen, chlorine, and hydrogen ? How is a molecular formula determined ? Illustrate your answer. 13. What is a molecular equation ? Give two illustrations. How does it differ from an ordinary chemical equation ? Of what use are such equations ? 14. Define () valence, (b} monad, dyad, triad, tetrad, pentad, (c) univalent element, bivalent element, (d) saturated compound, (e) unsaturated compound. 15. What is the valence of hydrogen ? Why ? Of oxygen ? Why ? How may valence be found by inspecting a binary formula ? What is the valence of NH 4 and OH ? 1 6. Illustrate the ways valence may be represented. 17. Distinguish between structural and empirical formulas. 1 8. What is the valence of sodium in (a} sodium chloride, (b) so- dium nitrate (NaNO 3 ), (c} sodium sulphate (Na 2 SO 4 ), (d) sodium hydroxide (NaOH) ? 19. What is the valence of sulphur in (a) snlphur dioxide (SO 2 ), (b) sulphur trioxide (SO 3 ), (c} hydrogen sulphide (H 2 S), (d) sulphuric acid, (e) copper sulphate (CuSO 4 ) ? (Suggestion. In oxygen acids, the oxygen valence balances the sum of the valence of the other elements.) 20. What is the valence of (#) aluminium in aluminium oxide (Al 2 Oo), (b} carbon in carbon tetrachloride (CC1 4 ), (c) phosphorus in phosphorus pentoxide (P 2 O.) ? I 180 Descriptive Chemistry. 21. What is the valence of (a} silver and chlorine in silver chloride (AgCl), () calcium and chlorine in calcium chloride (CaCl 2 ), (<;) oxy- gen in water, (d) oxygen and calcium in calcium oxide or lime (CaO) ? PROBLEMS. 1. The vapor densities of certain gases is as follows : (#) hydro- chloric acid 18.25, (b} chlorine 35.5, (c) ammonia 8.5, (d) nitrogen 14, 0) steam 9. Calculate the molecular weight of each. 2. Calculate the simplest formula of the compounds which have the indicated composition: (a} N = 82.353, H = 17.647; () O = 30, Fe (iron) =70; (c) H = i, C = 12, K (potassium) =39, O = 48. 3. A liter of sulphurous oxide gas (SO 2 ) weighs 2.8672 gm. What is the molecular weight of this compound ? 4. If 1500 cc. of carbon monoxide gas (CO) weigh 1.8816 gm., what is the molecular weight of the compound ? 5. Calculate the molecular formula of the compounds corresponding to the following data: (a) C = 73.8, H = 8.7, N = 17.1, vapor density = 80.2; () C=92.3, H = 7.7, vapor density =38.8 ; (c) C = 39.9, H = 6.7, O =53.4, vapor density = 30.5. 6. What volumes of factors and products are represented by the equations (a) H 2 + C1 2 = 2 HC1, () 2 H, + O 2 = 2 H,O, (c) 3 H, + N 2 = 2 NH 3 , (d) N 2 + O 2 = 2 NO, (e) 2 NO + O, = 2 NO, ? 7. If 20 1. of hydrogen are allowed to interact with 10 1. of chlo- rine, (a) how many liters of hydrochloric acid gas are produced, and () which gas and how much remains ? 8. How many liters of hydrogen gas can be obtained from 4 1. of hydrochloric acid gas ? 9. If 91.462 gm. of silver, when heated in chlorine, yield 121.4993 gm. of silver chloride, what is the atomic weight of chlorine ? (Assume Ag = 108.) 10. How many liters of the component gases can be obtained by the decomposition of 6 1. of ammonia gas ? 11. Find the simplest formulas of the substances having the follow- ing composition : (a) H = 1.58, N = 22.22, O = 76.19 ; (^) O = 47.52, N = 13.86, K = 38.61. 12. A certain weight of copper oxide, when heated in a current of hydrogen, lost 59.789 gm. of oxygen and formed 67.282 gm. of water. (a) If O = 16, what is the atomic weight of hydrogen ? () If H = i, what is the atomic weight of oxygen ? CHAPTER XIV. CARBON AND ITS OXIDES CYANOGEN. Occurrence of Carbon. Uncombined carbon is found pure in nature as diamond and graphite ; in a more or less impure state it occurs as coal and similar substances, which are included in the term amorphous carbon. Car- bon forms a vast number of compounds, natural and artificial. Combined with hydrogen and oxygen, and occasionally with nitrogen also, it is an essential constitu- ent of plants and animals. Meat, starch, fat, sugar, wood, cotton, paper, soap, wool, wax, flour, albumen, and bone contain carbon. It is also a component of carbon dioxide and of carbonates, such as limestone, chalk, and marble. Illuminating gases, kerosene and other products of petro- leum, turpentine, alcohol, chloroform, ether, and similar liquids are compounds of carbon. It is estimated that 0.22 per cent of the weight of the earth's crust is carbon. Diamond is pure crystallized carbon. It is found in only a few places in the earth. When taken from the mine, diamonds are rough-looking stones ; some are crystals, some are rounded like peas, and many are irregular ; they must be cut and polished to bring out the luster and make them sparkle (Fig. 24). The highly prized diamonds are colorless and without a flaw, and are said to be "of the first water " ; yellow ones from South Africa are common, and occasionally a blue, pink, red, or green one is found ; a very impure variety is black. 181 182 Descriptive Chemistry. The diamond is insoluble in all liquids at the ordinary temperature, has the high specific gravity of 3.5, and is the hardest known substance. It is brittle and may be shattered by a blow with a hammer. Crystal. Rough. FIG. 24. Diamonds. Cut. Diamonds have always been prized as gems on account of their beauty, rarity, and permanency. Besides being worn as jewels, they are used to cut glass, and the powder and splinters (known as bort) are used to grind and polish diamonds and other hard gems. The im- pure variety which comes from Brazil, and is called carbonado, is set into the end of the " diamond drill,' 1 which is used extensively for boring artesian wells and drilling hard rocks. The diamond was formerly found in gravel deposits in India, and in later years in Brazil. Since 1867, however, about 95 per cent of the dia- monds of commerce have come from South Africa. They occur in a bluish volcanic rock along the Vaal River, and especially near Kimberley. Over eight tons of diamonds have been found in South Africa in the last twenty-five years ! The successive investigations of Lavoisier, Dumas, and Davy, ex- tending from 1772 to 1814, showed that diamond is carbon, for when pure diamond was burned in oxygen, the only product was carbon dioxide. This result, which ad- mits of no doubt, has been verified by many famous investigators. Diamonds have been made by Moissan. He dissolved pure char- coal in melted iron, and poured the molten mass into water. The sur- face was so suddenly cooled that a tremendous pressure was exerted FlG. 25. Artificial diamonds (enlarged) prepared by Moissan. Carbon and its Oxides. 183 by the expanding iron inside the crust. This pressure caused the cool- ing carbon to crystallize into diamond. The crystals were very small, most of them were black, a few were white, but all had the properties of the diamond (Fig. 25). Large diamonds have a fascinating history, since most of them have passed through many hands before finding a place among royal jewels. The largest is the Orloff, which weighs 194! carats, and is in the scepter of the Czar of Russia. 1 The Kohinoor, which now weighs about 106 carats, is one of the crown jewels of England. Graphite is a soft, black, shiny solid, which is smooth and soapy to the touch. Pure graphite is carbon. It occurs native in large quantities and in many places. One va- riety is found in abundance at Ticonderoga, New York. Other famous localities are Ceylon, eastern Siberia, Bava- ria, and Italy. Sometimes crystals and grains are found, but it usually occurs in flaky masses or slabs. Unlike diamond, graphite is a good conductor of electricity and is often used to coat moulds in electrotyping. It is so soft that it blackens the fingers and leaves a black mark on paper when drawn across it. This property is indicated by the name graphite, which is derived from a Greek word (grap/iein) meaning to write. It resembles diamond in its insolubility in liquids at the ordinary temperature. Its specific gravity is 2.2, being considerably lighter than dia- mond. It produces only carbon dioxide when burned in oxygen ; but unlike diamond, it turns into carbon dioxide by heating to a very high temperature in the air. Graphite was once supposed to contain lead, and rs even now often incorrectly called " black lead " and plumbago. It is used to make stove polish and protective paints, as a lubricant where oil cannot be used, as' the principal ingredient of 1 A carat equals 3J Troy grains (or 0.205 gm.). The term is derived from the carob bean, which was used for ages by the diamond merchants of India as a small weight. 184 Descriptive Chemistry. graphite crucibles, in which metals are often melted, and in making electrodes for the huge electric furnaces. Immense quantities of graphite are consumed in the manufacture of lead pencils. The graphite is washed free from impurities, ground to a fine powder, mixed with more or less clay, and then pressed through perforated plates, from which the "lead" issues in tiny rods. These are dried, cut into the proper lengths, baked to remove all traces of moisture, and then inserted in the wooden case. In the United States in 1902 over four million pounds of graphite were mined, and over thirty-two million pounds were imported. Molten iron and other metals dissolve carbon, and when the metals cool the carbon crystallizes as graphite. Moissan incidentally obtained considerable graphite in making diamonds. Artificial graphite is now a commercial article (see Chapter X). Amorphous Carbon is a broad term, including all vari- eties of coal and charcoal, lampblack, and gas carbon. They are the non-crystalline forms of impure carbon. The word amorphous means literally "without form," and it is often used to designate soft, powdery, and uncrys- tallized substances. Coal is a term applied to several varieties of impure carbon. It may be regarded as the final product derived from vegetable matter by heat and pressure to which it was subjected through long geological periods. Ages ago the vegetation was exceedingly dense and luxuriant upon land slightly raised above the sea. In process of time this vegeta- FiG. 26. Section of part of the earth's crust near Mauch Chunk, Penn., showing layers of coal. tion decayed, accumulated, and slowly became covered with sand, mud, and water. The heat of the earth and the enormous pressure of the overlaying deposits changed the vegetable matter into more or .less Carbon and its Oxides. 185 impure carbon. This series of geological and chemical changes was repeated, and as a result we find in the earth layers or seams of carbo- naceous matter varying in thickness and composition (Fig. 26). These are the coal beds. Coal .beds contain proofs of their vegetable origin, viz., impressions of vines, stems, and leaves of plants, and similar vegetable substances FIG. 27. Fossil found in a FIG. 28. Section of coal as seen through coal bed. a microscope. (Fig. 27). A thin section of coal examined through a microscope re- veals a distinct vegetable structure (Fig. 28). There are three principal kinds of coal, (i) Bitumi- nous or soft coal is used to make illuminating gas, coke, and as a fuel for steam ; it burns with a smoky flame, and in burning produces much volatile matter. (2) Anthra- cite coal is hard and lustrous. It ignites with difficulty, burns with little or no flame, and produces an intense heat. It is used mainly for domestic purposes, heating and cooking, especially in eastern United States. (3) Lig- nite or brown coal is the least valuable as fuel. It often shows the woody fiber and was probably formed much later than the other varieties. Peat, strictly speaking, is not coal, though it is used as fuel in some places, espe- cially in Ireland and Holland. It is formed by the slow i86 Descriptive Chemistry. decay of roots and other vegetable matter under water, and represents an early stage of coal formation. The average composition of different kinds of coal is seen by the following table : , KIND. CARBON. VOLATILE MATTER. ASH. WATER. Lignite .... TO Q 2O Q jO 2 18 -> w -y Bituminous .... 74-53 I5-I3 10.34 Anthracite .... 91.64 6.89 1.47 Some anthracite coals contain as much as 95 to 99 per cent of carbon, and some bituminous coals as little as 65 per cent. Peat and wood contain still less carbon, but FIG. 29. Coal fields in the United States. more volatile matter. The volatile matter includes nitro- gen, hydrogen, and sulphur. These facts show that vege- table matter, in passing through the changes which finally Carbon and its Oxides. end in coal, loses volatile matter, Anthracite coal, which is found at different depths and associated with rocks of different ages, shows that it was formed from the bitumi- nous variety by the great pressure caused by mountain building. Hence it loses volatile matter and becomes hard. Coal is widely distributed in the crust of the earth, but the deposits vary in extent and quality. It underlies about one sixth of the area of the United States, the anthracite variety covering less than five hundred square miles in eastern Pennsylvania (Fig. 29). The United States now leads the world in coal production, furnishing about one third of the total supply. England for many years headed the list, and even now furnishes a large amount, for its deposits are extensive (Fig. 30). Charcoal is a variety of amor- phous carbon obtained by heating wood, bones, ivory, and other organic matter in closed vessels, or by partially burning them in the air. Th? process consists essentially in driving off the vola- tile matter and retaining the carbon. Wood Charcoal is a black, brittle solid, and often has the form of the wood from which it is made. It is insoluble, though its mineral impurities may be removed by acids. It burns without lame or much smoke, and leaves a white ash. The compact varieties conduct heat and electricity, but porous charcoal is a poor conductor. It resists the action of many chemicals; hence fence posts, telegraph poles, and wooden piles are often charred before being JKITISH COALFIELDS FlG. 30. Coal deposits in the British Isles. 1 88 Descriptive Chemistry. put into the ground. Most varieties are very porous, and when thrown upon water charcoal floats, owing to the presence of air in its pores. Its porosity makes charcoal an excellent absorber of gases, some varieties absorbing ninety times their bulk of ammonia gas. Sewers and foul places are sometimes purified by charcoal. It will also absorb colored substances from solutions. This is espe- cially true of animal charcoal (see below). Foul air and water may be partially purified by charcoal, which forms the essential part of many water filters in houses. Char- coal used for such a purpose, however, must be renewed or often heated to redness; otherwise it becomes clogged and contaminated. Charcoal is never pure carbon, the degree of purity depending upon the kind of wood used, as well as the temperature and method employed. Besides the uses of charcoal mentioned above, it is used as a fuel, in the manufacture of steel and of gunpowder, and as a medicine. It reduces oxides when heated with them, thus 2 CuO + C = 2 Cu + CO 2 Copper Oxide Carbon Copper Carbon Dioxide Wood charcoal is made either in a charcoal pit or kiln, or in a large retort. Where wood is plentiful, it is loosely piled into the shape shown in Figure 31, and covered with turf to prevent free access of air, though small holes are left at the bottom and a larger one at the top of a central flue, so that sufficient air can pass through the pile. The wood is lighted, and as it slowly burns care is taken to regulate the supply of air, so that the wood will smolder but not burn up. The volatile matter escapes and charcoal remains, the average yield being about 20 per cent of the weight of the wood. This method is crude, uncertain, and wasteful. Much charcoal is now made by heating wood in closed retorts, no air whatever being admitted. By this method, which is called dry or destructive distillation, the yield of charcoal is 30 per cent and all the volatile matter is saved. In the Carbon and its Oxides. 189 ordinary combustion of wood, the hydrogen forms water and the oxy- gen forms carbon dioxide ; but in dry distillation, where no oxygen is present, much of the hydrogen forms volatile compounds with the car- bon and oxygen. Among these volatile products are methyl alcohol FIG. 31. Wood arranged for burning into charcoal. and acetic acid. These are commercial substances, and contribute to the profit of the process. More or less charcoal is obtained by heating any compound of carbon, e.g. sugar or starch, the charring being a test for carbon. Animal Charcoal or Bone Black is made by heating bones in a closed vessel, and by heating a mixture of blood and sodium carbonate. It contains only about 10 per cent of carbon, but this carbon is dis- tributed throughout the porous mineral matter of the bone, which is almost entirely calcium phosphate. Under the name of ivory black, animal charcoal is used as a pigment, especially in making shoe-black- ing. It is extensively used to remove the color from sugar sirups, oils, and other liquids colored by organic matter. Coke is made by expelling the volatile matter from soft coal, somewhat as charcoal is made from wood. It is left in the retorts when coal is distilled in the manufacture of illuminating gas. On a large scale it is made by heating a special grade of soft coal in huge brick ovens, shaped like a beehive, from which air is excluded after combus- tion begins. Sometimes the coke is made in closed retorts constructed so as to save the by-products, ammonia, tar, 190 Descriptive Chemistry. organic compounds, and combustible gases. This method not only yields more coke, but is also more profitable be- cause the by-products are sold and the combustible gas is used to heat the retorts. Coke is a grayish, porous solid, harder and heavier than charcoal. It burns with no smoke and a feeble flame. It contains about 90 per cent of car- bon, the rest being the mineral matter originally in the coal. Immense quantities of coke are used in the manufacture of iron and steel. It is superior to coal for this purpose, because it gives a greater heat when burned, reduce's oxides easily, and contains little or no sulphur or other substances harmful in the iron industries. Coke is the fuel used in making nine tenths of the pig iron in the United States, and over twelve million tons (or about three fourths of the total amount) are made annually in the Connellsville district, near Pittsburg, Pennsylvania. Gas Carbon is amorphous carbon which is gradually deposited upon the inside of the retorts used in the manufacture of illuminating gas. It is a black, heavy, hard solid, and is almost pure carbon. It is a good conductor of electricity, and is extensively used for the manufacture of the carbon rods of electric lights and for plates of electric batteries. Lampblack is prepared by burning oil or oily substances rich in carbon in a limited supply of air. The dense smoke, which is mainly finely divided carbon, is passed through a series of condensing cham- bers, where it is collected upon coarse cloth or a cold surface. Its formation is illustrated on a small scale by a smoking lamp, and the soot deposited is the same as lampblack. Lampblack is one of the purest forms of amorphous carbon, and it is used in making printer's ink and certain black paints. Allotropism. Diamond, graphite, and amorphous car- bon, though exhibiting essentially different properties, are identical in composition. All are carbon. They can be changed into one another, the amorphous form into graph- ite and finally into diamond and the diamond into amor- phous carbon. Each burns in oxygen and the product is carbon dioxide. Furthermore, the same weight of each Carbon and its Oxides. 191 forms the same weight of carbon dioxide, i.e. when 12 gm. of each are burned, 44 gm. of carbon dioxide are always produced. There is no doubt about their identity, though no one has explained it. The property of assum- ing more than one elementary form is called allotropism or allotropy (from Greek words meaning another form). The more uncommon form is called an allotrope or an allotropic modification of the other. It is believed by some that allotropism is due to a difference in the number of atoms in a molecule of the element. OXIDES OF CARBON. Carbon and Oxygen do not unite at the ordinary tem- perature. But when carbon is heated in air, in oxygen, or with some oxides, carbon dioxide (CO 2 ) is formed ; if the supply of oxygen is limited, then carbon monoxide (CO) is formed. Occurrence and Formation of Carbon Dioxide. The occurrence of carbon dioxide in the atmosphere and in many natural waters has already been mentioned. It is the main product of ordinary combustion, respiration of animals, and decay. In all these processes the carbon comes from organic matter, while the oxygen comes from the air, from the organic matter, or from both. Ordinary combustion is a chemical combining of carbon and oxygen. Hence, when carbon or a substance contain^ ing it is burned, carbon dioxide is formed. The equation for this change is C + 2 C0 2 Carbon Oxygen Carbon Dioxide Carbon dioxide is formed by the combustion of such com- mon substances as wood, coal, charcoal, coke, oils, waxes, 192 Descriptive Chemistry. cotton, bone, starch, sugar, meat, bread, alcohol, camphor, and illuminating gas. The continuous oxidation of the tissues and foods in the body produces carbon dioxide (see Relation of Oxygen to Life). And if we exhale the breath through a glass tube into limewater, the carbon dioxide which is in the breath turns the limewater milky the usual test for carbon dioxide. The equation for the change is CO 2 + Ca(OH) 2 = CaCO 3 + H 2 O Carbon Dioxide Limewater Calcium Carbonate When vegetable and animal matter decays, carbon dioxide is formed. Many kinds of organic matter fer- ment, especially those containing sugar. By alcoholic fermentation the sugar changes into carbon dioxide and alcohol (see Alcohol), thus C 6 H 126 = 2CO 2 + 2C 2 H 6 O Sugar Carbon Dioxide Alcohol The Preparation of Carbon Dioxide is usually accom- plished by the interaction of a carbonate and an acid. Calcium carbonate (limestone or marble) and hydrochloric acid are usually used. The operation may be easily per- formed in any glass vessel by pouring the acid upon the carbonate. The equation for the chemical change is CaCOg + 2HC1 = CO 2 + CaCl 2 + H 2 O Calcium Carbon Calcium Carbonate Dioxide Chloride This gas may also be prepared by heating matter con- taining carbon, or by strongly heating carbonates (as in making lime), thus CaCOg CO 2 + CaO Calcium Carbonate Carbon Dioxide Lime Carbon and its Oxides. Properties of Carbon Dioxide. This gas has many important properties besides those mentioned under The Atmosphere. It has a slight taste and odor, but no color. It is one and a half times heavier than air, and a liter under standard conditions weighs 1.977 gm. On ac- count of its weight it can be collected by downward dis- placement and poured from one vessel to another. For the same reason, it is often found at the bottom of old or deep wells, in some valleys near lime kilns or volcanoes, and in mines after explosions. At the ordinary tempera- ture and pressure, water dissolves its own volume of carbon dioxide. Under increased pressure more gas dis- solves, which escapes readily when the pressure is re- moved. Hence " soda water," which is made by forcing carbon dioxide into water, effervesces and froths when drawn from the soda fountain. Many natural waters and manufactured beverages (such as champagne and beer) sparkle and effervesce for the same reason. This gas may be liquefied by subjecting it to high pressure and low temperature. It was first liquefied by Faraday by the method used for chlorine. Liquid carbon dioxide is now made in large quantities by forcing the gas into steel cylinders by powerful pumps, the gas being obtained in many cases from the fermenting vats of breweries. When a cylinder of liquid .carbon dioxide is opened, the liquid evaporates so rapidly that a portion of it becomes a white, snowlike solid. Both the liquid and solid carbon dioxide are articles of commerce, and are sometimes used to prepare "soda water," to extinguish fires, to improve wines, and to produce very low temperatures. Carbon dioxide extinguishes burning objects, such as a blazing stick or lighted candle; indeed, air containing from 2.5 to 4 per cent of carbon dioxide will extinguish 194 Descriptive Chemistry. small flames. Hence the gas is often used to extinguish fires. Many small fire extinguishers contain sodium carbonate and sulphuric acid, so arranged that when desired, carbon dioxide gas may be generated from them under pressure. A stream of the gas forced upon a small blaze will often prevent a serious fire. In other forms, the carbon dioxide, which is similarly generated, forces water from the extinguisher. Relation of Carbon Dioxide to Life. Animals die when put into carbon dioxide. It cuts off the supply of oxygen as water does from a drowning man. The presence of a small quantity in the air is objectionable, since it is said to produce headache and drowsiness; but much of the dis- comfort felt in badly ventilated rooms and attributed to carbon dioxide is doubtless due to water vapor, and to poisonous substances produced from the organic mat- ter exhaled from the lungs. On the other hand, carbon dioxide is an essential food of plants. Through their leaves and other green parts they absorb carbon dioxide from the atmosphere, decompose it, reject the oxygen, and store up the carbon in the form of starch. The sunlight and the green coloring matter aid the plant in manufac- turing its food out of the water (obtained through the roots from the soil) and the carbon of the carbon dioxide ob- tained from air. Plants thus serve to keep the atmosphere free from an excess of carbon dioxide, the proportion present in the air being very small and practically con- stant. Carbonic Acid. Carbon dioxide gas is often called carbonic acid gas, or simply carbonic acid. It is believed that carbon dioxide, when passed into water, combines with the water and forms a weak, unstable acid, which is, strictly speaking, carbonic acid. The equation for this change is Carbon and its Oxides. 195 CO 2 + H 2 = H 2 CO 3 Carbon Dioxide Carbonic Acid Such a solution reddens blue litmus and decolorizes pink phenolphthal- ein. Carbonic acid has never been obtained free, and is so unstable that it easily breaks up by gentle heat into carbon dioxide and water, thus H,CO 3 = CO 2 + H 2 0. Carbon dioxide is sometimes called carbonic anhydride, to denote its relation to the acid. Carbonates are salts corresponding to the unstable carbonic acid. They are stable compounds. The most abundant natural carbonates are those of calcium, magne- sium, and iron. Immense quantities of sodium and potas- sium carbonates are manufactured. A few carbonates are formed by direct combination of an oxide and carbon dioxide, but most of them are formed by passing carbon dioxide into the corresponding hydroxide, thus CO 2 + Ca(OH) 2 CaCO 3 + H 2 O Calcium Hydroxide Calcium Carbonate Many carbonates are insoluble in water, e.g. calcium carbonate, the test for carbon dioxide depending upon this fact. Others, e.g. sodium and potassium carbonate, are very soluble. There are two classes of carbonates, the normal and the acid. Normal sodium carbonate is Na 2 CO 3 , and acid sodium carbonate is HNaCO 3 . The latter is often called sodium bicarbonate. Normal calcium carbonate is CaCO 3 , and acid calcium carbonate is H 2 Ca(CO 3 ) 2 ; 4he latter is unstable, and is easily decomposed by heat into normal calcium carbonate. Composition of Carbon Dioxide. If a known weight of pure car- bon, such as diamond or graphite, is burned in oxygen, it is found that for 12 parts of carbon used there are 44 parts of carbon dioxide formed. Hence 12 parts of carbon unite with 32 parts of oxygen. The vapor density of the gas is 22, and the molecular weight must be 44. These facts necessitate the formula CO 2 . 196 Descriptive Chemistry. History of Carbon Dioxide. This gas was described in the seven- teenth century by Van Helmont, who called it gas sylvestre. He prepared it by the interaction of acids and carbonates, detected it in mineral water, and observed its formation during combustion and fer- mentation, as well as its action on animals and flames. Black, in 1755, showed that carbon dioxide is essentially different from ordinary air and that the gas is readily obtained from magnesium and calcium carbonates. Since the gas was combined or " fixed " in these substances, he called the gas fixed air. His work was verified in 1774 by Bergman, who called the gas acid of air. Lavoisier first proved it to be an oxide of carbon. Carbon Monoxide is formed when carbon is burned in a limited supply of air, thus C + O CO Carbon Oxygen Carbon Monoxide If carbon dioxide is passed over heated charcoal, the prod- uct is carbon monoxide. That is, carbon reduces carbon dioxide to carbon monoxide, the equation for the change being C0 2 + C 2 CO Carbon Monoxide This chemical change takes place in every coal fire. The oxygen of the air entering the bottom of the fire unites with the carbon to form carbon dioxide ; the latter gas in passing through the hot carbon of the fire is reduced to carbon monoxide. Some of the carbon monoxide escapes and some burns with a flickering bluish flame on the top of the fire. If steam is passed over red-hot coke or charcoal, a mixture of carbon monoxide and hydrogen is produced. This mixture enriched by vapor from oils is known as water gas (see Water Gas) . Carbon monoxide is usually prepared by gently heating a mixture of oxalic acid and sulphuric acid in a flask, and Carbon and its Oxides. 197 collecting the gaseous product over water. The oxalic acid decomposes thus C 2 H 2 O 4 = CO + CO 2 + H 2 O Oxalic Acid Carbon Monoxide Carbon Dioxide The carbon dioxide may be removed by passing the mixed gases through a solution of sodium hydroxide. Carbon monoxide is a gas without color, odor, or taste, and is only slightly soluble in water. It burns with a bluish flame, forming carbon dioxide, thus 2 CO + .O 2 2CO 2 Carbon Monoxide Carbon Dioxide Carbon monoxide is extremely poisonous, and it is doubly dangerous because its lack of odor prevents its detection in time to escape its stupefying effect. Many deaths have been caused by breathing air containing it. Carbon mo- noxide forms a compound with one of the constituents of the blood, and those who have been poisoned by it cannot be revived by air, as in the case of suffocation by carbon dioxide. It is a constituent of ordinary illuminating gas, and care should always be taken to prevent the escape of illuminating gas (as well as the gas from a coal stove or furnace) into rooms occupied by human beings. At a high temperature carbon monoxide unites easily with oxygen, and is, therefore, an important agent in the reduction of iron ores in the blast furnace. This action might be rep- resented thus Fe 2 3 + 3 CO = 2Fe + 3 CO 2 Iron Oxide Carbon Monoxide Iron Carbon Dioxide Carbon monoxide, which is sometimes called carbonic oxide, forms no acid and therefore no salts. It does not make limewater milky, thus being readily distinguished from carbon dioxide. Its blue flame dis- 198 Descriptive Chemistry. tinguishes it from all other gases which burn. It unites directly with chlorine to form carbonyl chloride (phosgene, COC1 2 ), and with some metals, forming metallic carbonyls, e.g. nickel carbonyl (Ni(CO) 4 ). Cyanogen is a compound of carbon and nitrogen having the composition corresponding to the formula (CN) 2 . It is a colorless gas, has the odor of peach kernels, is exceed- ingly poisonous, and burns with a purplish flame. It may be prepared by heating mercuric cyanide (Hg(CN) 2 ). Cyanogen is a radical, and in compounds it acts like an element. Its corresponding acid is hydrocyanic or prus- sic acid (HCN). This acid is prepared by heating a cyanide with sulphuric acid, "just as hydrochloric acid is obtained from a chloride. The solution smells like peach kernels, and is one of the most deadly of all known poisons. Potassium cyanide is a white, deliquescent solid. It is a deadly poison. Large quantities are used in gold and silver plating and in the " cyanide process " of extracting gold from its ores, as described under that metal. Other cyanogen compounds are cyanic acid (CNOH), sulpho- cyanic acid (CNSH), and potassium sulphocyanate (CNSK). The last is a white, crystallized salt, which produces a beautiful red solution when added to certain soluble iron compounds, and is therefore used to detect this metal. Salts of complex acids related to hydrocyanic acid are used in dyeing, many being prepared from the most common one potassium ferrocyanide or yellow prussiate of potash. They will be described in the chap- ter on Iron. EXERCISES. 1. What is the symbol and atomic weight of carbon? 2. In what forms does free carbon occur in nature? Name ten famil- iar solids, three liquids, and two gases containing carbon. What pro- portion of the earth's crust is carbon? Carbon and its Oxides. 199 3. What is diamond? How could the correctness of your answer be shown? State (a} the source, (b) the properties, and (c) the uses of diamonds. Give a brief account of one or more famous diamonds. 4. What is graphite? What is its chemical relation to diamond, and how could this relation be proved ? State (a) the source, (b} the properties, and (c) the uses of native graphite. 5. What is {a} black lead, () plumbago, (c) bort, (d) carbonado, (e) native graphite, (/) artificial graphite? 6. Give a brief account of the manufacture of lead pencils. What is the literal meaning of graphite? 7. Review artificial graphite (see Chapter X). 8. What does the term amorphous carbon include? Does the car- bon in these impure forms differ chemically from diamond and graphite? 9. How was coal formed? Give several proofs of its origin. State the properties and uses of (a) bituminous coal, (b} anthracite coal, and (V) lignite. What besides carbon does it contain? Where is coal found ? 10. W T hat is charcoal ? State (a} the properties, and (b} the uses of wood charcoal. Give a brief account of both methods of preparing wood charcoal. State the preparation, properties, and uses of animal charcoal. u. What is coke? How is it made? What are its properties? How is it related to the iron industries? 12. What is gas carbon? What is its source? State its properties and uses. 13. What is lampblack? State its method of preparation, properties, and uses. 14. Define and illustrate (a} amorphous, and (b} allotropism. 15. Develop the topics: (a} carbon is a reducing agent, (b} carbon monoxide is a reducing agent, (c) diamond, graphite, and pure amor- phous carbon illustrate allotropism. 1 6. What is (a) hard coal, (b) soft coal, (c) peat, (d} boneblack, (e) soot, (/) lampblack, (g) lignite, (h) electric light carbon? 17. Give the names and formulas of the two oxides of carbon. How is each formed from carbon and oxygen? 1 8. Describe the occurrence and formation of carbon dioxide. What is always obtained by burning a substance containing carbon ? Give the simplest equation for this chemical change. 19. Describe fully the action of carbon dioxide on limewater. Give the equation for the reaction. 2OO Descriptive Chemistry. 20. What is the relation of carbon dioxide to (a} respiration, (b) fer- mentation of sugar, (c) decay, (d) making lime ? 21. What is the test for (a) carbon, () carbon monoxide, (c} car- bon dioxide? 22. Describe the usual method of preparing carbon dioxide. Give the equation for the reaction. State its properties. 23. Describe liquid and solid carbon dioxide. How are they pre- pared ? For what are they used ? 24. What is the relation of carbon dioxide to animal and to plant life? 25. State fully the relation of carbon dioxide to the unstable acid H.,CO 3 . Give the equations for the formation and decomposition of this acid. 26. What are carbonates? Name three. How are they formed? What are their properties ? 27. What is (a) " soda water," () carbonated water, (c) carbonic acid, (d) carbonic oxide, (e) carbonic anhydride, (/) limestone or marble? 28. What is the difference between (a} sodium carbonate and sodium bicarbonate, and (b) calcium carbonate and acid calcium carbonate ? 29. Why is (a) CO 2 the formula of carbon dioxide, and (b} CO of carbon monoxide? 30. State briefly the history of carbon dioxide. 31. Give a brief account of (a} Black, (b) Van Helmont, and (c) Bergman. 32. Illustrate the law of multiple proportions by the oxides of carbon. 33. Give the equations for (a} the oxidation of carbon to carbon monoxide, (^) the reduction of carbon dioxide to carbon* rnonoA^-.. 34. How is carbon monoxide (#) formed, and (<) usually prepared ? 35. What is the relation of carbon monoxide to water gas? 36. What are the properties of carbon monoxide? 37. Illustrate Gay-Lussac's law by the combustion of carbon mo- noxide (2 CO + O., = 2 CO 2 ) . 38. Illuminating gas, water gas, and the gas which escapes from a coal fire are poisonous. Why? 39. What is cyanogen? Hydrocyanic acid? Describe potassium cyanide. For what is it used? Describe ootassium sulphocyanate. State its chief use. Carbon and its Oxides. 101 40. The specific gravity of charcoal is about 1.5. Why does it float on water? 41. How can carbon monoxide and carbon dioxide be changed into each other? 42. Review (a) combustion, () solution of gases (especially carbon dioxide) in water, (c} respiration. 43. State and explain the various chemical changes which occur from the entrance of oxygen (in the air) below the grate of a red-hot coal fire to the end of the burning of the carbon monoxide at the top of the coal. PROBLEMS. 1 . How many grams of calcium carbonate are needed to prepare 132 gm. of carbon dioxide ? 2. What weight of carbon burned in air will produce n gm. of carbon dioxide ? 3. Calculate the percentage composition of (a} calcium carbonate, () carbon monoxide, (c) carbon dioxide, (d) magnesium carbonate. 4. What per cent of carbon (by weight) is contained in carbon monoxide and in carbon dioxide ? 5. If 20 gm. of carbon are heated in the presence of 44 gm. of carbon dioxide, (a) what weight of carbon monoxide is formed, and (<) what weight, if any, of carbon remains ? 6. How many liters of carbon dioxide must be passed over red-hot charcoal to yield 84 gm. of carbon monoxide ? 7. How much carbon dioxide () by weight and () by volume is in the air of a room 6 m. long, 4 m. wide, and 3 m. high, if there is i vol. of carbon dioxide in 1000 vol. of air ? 8. What weight of water must be decomposed to furnish enough oxygen to form (with pure carbon) 44 gm. of carbon dioxide ? 9. How many grams of calcium carbonate will produce 15 1. of carbon dioxide ? 10. If a piece of pure graphite weighing 7 gm. is burned in oxygen, what volume of carbon dioxide is formed ? CHAPTER XV. HYDROCARBONS METHANE ETHYLENE ACETYLENE -ILLUMINATING GAS FLAME BUNS EN BURNER - OXIDIZING AND REDUCING FLAMES. Hydrocarbons are compounds of carbon and hydrogen. They number about two hundred, and their properties vary between wide limits. They are found in petroleum and its products (kerosene, naphtha, lubricating oils, par- affin wax, etc.), in coal tar, in coal gas and natural gas, and in some essential oils, such as turpentine. On a large scale they are prepared by the destructive distillation of petroleum, wood, coal, and coal tar. Indirectly the hydro- carbons are the source of many other compounds of car- bon, which are extensively used in numerous industries. The existence of so many hydrocarbons is due to the fact that atoms of carbon have power to unite with themselves. This property gives rise to compounds which form natural groups or series. Simple rela- tions exist between many hydrocarbons, especially between members of the same series. The consecutive members of a series differ in com- position by CH 2 . Thus, in the methane series, methane is CH 4 and ethane is C 2 H 6 ; in the ethylene series, ethylene is C 2 H 4 and propylene is C 3 H 6 ; in the acetylene series, acetylene is C 2 H 2 and allylene is C 3 H 4 ; and in the benzene series, benzene is C 6 H G and toluene is C-H 8 . These series are called homologous series. Methane is found in coal mines, being a gaseous prod- uct of the processes which changed vegetable matter into coal. It is called fire damp by miners. It is also formed in marshy places by the decay of vegetable matter under water, and is therefore often called marsh gas. Methane. 203 It is a constituent of natural gas and petroleum, and forms a large proportion of the illuminating gas obtained by heating coal. Methane is usually prepared in the laboratory by heating a mixture of sodium acetate, sodium hydroxide, and quicklime in a hard glass or metal vessel, and collecting the gaseous product over water. It may also be prepared by the interaction of aluminium carbide and water, thus A1 3 C 4 +i2H 2 0= 3CH 4 + 4A1(OH) 8 Aluminium Carbide Water Methane Aluminium Hydroxide Methane has no color, taste, or odor. It burns with a pale, luminous flame. A mixture of methane with oxygen or air explodes violently when ignited by a spark or flame. Terrible disasters occur in coal mines from this cause. The products of the explosion are carbon dioxide and water, thus- CR4 + 2 Q 2 = co? + 2 H 2 Methane Oxygen Carbon Dioxide Water The carbon dioxide, called choke damp or black damp by the miners, often suffocates those who escape from the explosion. Other members of the methane series are ethane (C 2 H 6 ), propane (C 3 H 8 ), butane (C 4 H 10 ). This series is also called the paraffin series, on account of the chemical indifference of its members. It has the general formula C n H 2n + 2- Butane and the succeeding fifteen or twenty members are liquids, and the highest members are solids. Chlorine and hydrocarbons interact, that is, chlorine replaces hydro- gen, atom for atom. Thus CH 4 + 2C1 = CH 3 Cr + HC1 Methane Chlormethane This chemical change is called substitution, and illustrates one of the methods used in preparing derivatives of carbon known as substitution products. The paraffins are saturated hydrocarbons. This means that the carbon in them is saturated, so to speak, with hydrogen, and has no tendency to unite directly with more atoms of hydrogen or other elements. 204 Descriptive Chemistry. Ethylene or olefiant gas is formed by the destructive distillation of wood and coal. It is usually prepared by heating a mixture of concentrated sulphuric acid and ethyl alcohol, and collecting the gas over water. The alcohol decomposes into ethylene and water, the latter being ab- sorbed by the sulphuric acid. The essential change is represented thus C^WgO = C^H^ -f- H^O Alcohol Ethylene Ethylene is a colorless gas, and has a pleasant odor. It can be condensed to a liquid, which by evaporation pro- duces a temperature as low as 140 C. It burns with a bright, yellow flame, and is one of the illuminating constit- uents of coal gas. When ethylene burns, the complete combustion is represented thus C 2 H 4 + 3 2 = 2C0 2 + 2H 2 Ethylene Carbon Dioxide Water If mixed with oxygen in this proportion and ignited, the mixture explodes. Other numbers of this series are propylene (C 3 H 6 ) and butylene (C 4 H 8 ). These are unsaturated hydrocarbons. Unlike the paraffins, they form addition products by uniting directly with other substances, especially chlorine, thus C 2 H 4 + C1 2 = C 2 H 4 C1 2 Ethylene Ethylene Chloride Ethylene chloride is one of the two dichlorethanes ; they have the same percentage composition, molecular weight, and formula (C 2 H 4 C1 2 ), but are very different compounds. They illustrate isomerism and are called isomers. This kind of isomerism is called metamerism. The difference in properties is believed to be due to a different arrangement of the atoms in the molecules. Isomerism occurs frequently among carbon compounds. Acetylene. 205 Acetylene is formed by the direct union of hydrogen and carbon when an electric arc is produced between two carbon rods in hydrogen gas. This method of formation, though not convenient, is interesting, because no other hy- drocarbon has as yet been directly built up from its elements. A small quantity is present in coal gas. It is also formed by the incomplete combustioft of coal gas, e.g. when the flame of a Bunsen burner strikes back and burns at the base (see Bunsen Burner). Acetylene is now prepared cheaply on a large scale by treating calcium carbide with water, thus CaC 2 + 2H 2 O = C 2 H 2 + Ca(OH) 2 Calcium Carbide Acetylene Acetylene is a colorless gas, and, if impure, has an offen- sive odor. It is poisonous if breathed in large quantities, but much less dangerous than gases containing carbon monoxide. It is lighter than air, its density being about 0.92. Water at the ordinary temperature dissolves its own volume of the gas. Reliable tests show that acetylene does not act upon any common metal or alloy, though it forms explosive compounds with salts of metals, especially copper. As a precaution, copper and brass are seldom used in large vessels containing or generating acetylene, though they might be safely used on small vessels like bicycle lamps. Under a pressure of 40 atmospheres and a temperature of 20 C. it liquefies. Cylinders of liquid acetylene have exploded, causing loss of life and destruction of property, and its use in this form has been pro- hibited in some localities. Under ordinary atmospheric conditions acety- lene will not explode. If compressed, it will explode when a spark or flame is brought near it. A mixture of acetylene and air, if ignited, explodes. The mixture to be explosive, however, must contain from about 3 to 65 per cent of acetylene (a condition hardly possible 206 Descriptive Chemistry. except from sheer carelessness), because the disagreeable odor reveals the presence of the gas. Acetylene must be used with the same precau- tion as any other illuminating gas. Acetylene is found by analysis to contain only carbon and hydrogen combined in the ratio of 12 to I by weight. Its vapor density is 13. Therefore its molecular weight must be 26 and its formula C 2 H.,. Acetylene is an unsaturated hydrocarbon, and like ethylene combines directly with bromine, hydrogen, and other elements. When passed into silver or copper solutions, it forms explosive compounds called acetylides (e.g. Ag 2 C 2 and Cu 2 C 2 ). Heated to a high temperature, it changes into other hydrocarbons, one being benzene, thus 3 C 2 H 2 = C 6 H 6 Acetylene Benzene At a very high temperature (about 800 C.) it decomposes into carbon and hydrogen. The change of acetylene into benzene illustrates po- lymerism. Polymers have the same percentage composition, but different molecular weights (see Isomerism). Acetylene as an Illuminant. Acetylene burns in the air with a luminous, smoky flame. But when air is mixed with the gas as the latter issues from a small opening, the mixture burns with a brilliant, white flame, which does not smoke. It is grad- ually coming into use as an illuminant. The flame is almost like sunlight, hence by the acetylene flame most colors appear the same as in daylight. It is also adapted for taking photographs, since its action closely resembles that of the sun. It is a diffusive light, and the flame is much smaller than an ordinary gas flame of the same lighting power (Fig. 32). FlG. 32. Relative size of acetylene and illuminating gas flames giving the same amount of light. The acetylene (smaller) flame consumes only one tenth as much gas an hour as the illu- minating gas flame. (One half actual size.) Petroleum. 207 With a proper burner the combustion of acetylene is complete, and may be represented thus 2C 2 H 2 + 5O 2 = 4CO 2 + 2 H 2 O Acetylene Oxygen Carbon Dioxide Water In most acetylene burners the gas issues from two small holes drilled at an angle, so that the jets strike each other and produce a flat flame (Fig. 33). Other holes, properly located, permit air to be drawn in mechanically by the acetylene as it rushes through the burner. The open- ings for the mixture are so fine that FIG. 33. Acety- the flame cannot strike back and cause FlG 34 ._ Acety- lene flame. an explosion (Fig. 34). lene burner. Generation of Acetylene. The ease with which acetylene is gener- ated can be shown by putting a little water in a test tube and then drop- ping in small lumps of calcium carbide. The gas bubbles through the liquid ; after the action has proceeded long enough to expel the air, the acetylene may be lighted by holding a burning match at the mouth of the tube. On a larger scale, the gas can be generated by putting the calcium carbide into a flask provided with a dropping funnel and de- livery tube, and allowing water to drop slowly upon the carbide ; the gas thus generated can be collected in bottles over water. There are two classes of commercial generators. In one, water is added to the calcium carbide, but in the other the carbide drops into the water. The intense heat liberated when calcium carbide interacts with water de- composes acetylene ; hence, a generator to be effective and safe should be constructed so that this heat will be absorbed. The first class of generators is dangerous, except when a small quantity of gas is desired, as on the lecture table or in a bicycle lantern. -In the second class, a small amount of calcium carbide drops automatically into a large vol- ume of water as fast as the gas is needed, thus insuring a pure, cool gas, and eliminating the danger of an explosion. A pound of calcium carbide yields about five cubic feet of acetylene gas. Petroleum is the source of many useful hydrocarbons. It is an oily liquid obtained from the earth in many parts of the world. In the United States the chief localities are 208 Descriptive Chemistry. Ohio, New York, Pennsylvania, West Virginia, Kentucky, Indiana, Colorado, Texas, and California. The immense deposits in Russia are in the Baku district on the Caspian Sea. Some is also found in Canada, India, Japan, and Austria. Crude petroleum is a thick liquid, with an unpleasant odor. Its color varies from straw to greenish black, and most kinds are greenish in reflected light. It usually floats upon water. Its composition is complex, but all varieties are essentially mixtures of many hydrocarbons. Ameri- can oils contain chiefly members of the paraffin series. Some varieties contain compounds of nitrogen and of sulphur. In some localities the oil issues from the earth, but it is usually neces- sary to drill through rocks and insert a pipe into the porous rock containing oil. At first the oil often "shoots' 1 out of the well in tremendous volumes, owing to the pressure of the confined gas, but after a time a pump is needed to draw it to the surface. The oil is then forced by powerful pumps through large pipes to central points for storage or for delivery to refineries, which are often many miles from the oil well. This network of pipes in the eastern United States is over 25,000 miles long. Some crude petroleum is used in making water gas (see below), and as fuel on locomotives and steamships, but most of it is separated into various commercial products. This process, which also involves purification, is called re- fining. The petroleum is distilled in huge iron vessels, and the vapors are condensed as they pass through coiled pipes immersed in cold water. Certain products are ob- tained from the residue left in the still. The different distillates, which are collected in separate tanks, are further separated and purified by redistillation. The commercial products obtained from the first distillation are cymogene, rhigolene, gasolene, naphtha, benzine, and kerosene. These liquids are mixtures Natural Gas. 209 of several different hydrocarbons. They are widely used as solvents, fuels, and in making gas. Kerosene is the well-known illuminating oil. Being the most valu- able product from petroleum, it is very carefully freed from inflammable liquids and gases, which might cause an explosion, and from tarry matter and semi-solid hydrocarbons, which would clog the wicks of lamps. This is done by agitating it successively with sulphuric acid, sodium hydroxide, and water. Commercial kerosene must have a legal flashing point. This is "the temperature at which the oil gives off sufficient vapor to form a momentary flash when a small flame is brought near its surface." In most states the flashing point is 44 C. (or iiiF.). From the residuum left in the still after the first distillation many grades of lubricating oil, vaseline, and paraffin wax are obtained by further treatment. Mineral lubricating oils have largely replaced animal and vegetable oils. Vaseline finds extensive use as an ointment. Paraffin wax is used to make candles, to water-proof paper, to extract oils from plants and flowers, and as a coating for many substances, thereby producing a smooth surface or facilitating slow combustion (as in parlor matches). The final residue is coke. Hydrocarbons are often extracted from it, some is made into electric light carbons, and some is used as a fuel. This vast industry yields over two hundred different commercial products, many of them being indispensable to the comfort and conven- ience of mankind. In 1901 the United States produced over 69,000,000 barrels of crude petroleum. The Origin of Petroleum is doubtful. Some think it was produced by the decomposition or slow distillation of plants and animals. Recently it has been suggested that it resulted from the interaction of water and metallic carbides, especially iron carbide, at great depths. Natural Gas is a combustible gas, which issues from the earth in many places. Methane is the principal constituent of the mixture. It is used as a fuel for heating houses, generating steam, and manufacturing iron, steel, glass, brick, and pottery. In Ohio, Indiana, and other gas-producing regions of the United States, wells, like petroleum wells, are drilled for the escape of natural 2io Descriptive Chemistry. gas, which is distributed to consumers through pipes similar to those used for illuminating gas. Enormous quantities are consumed in the United States, the annual product being valued at over $20,000,000. Illuminating Gas. Besides acetylene there are other kinds of illuminating gas. Coal gas and water gas are the most common. Coal Gas is made by distilling bituminous coal and puri- fying the volatile product. The hydrogen in the coal passes off partly as free hydrogen, and partly in combina- tion with carbon as hydrocarbons, and with nitrogen as ammonia. The ammonia, carbon dioxide, and sulphur compounds are regarded as impurities, and are removed before the gas is sent to the consumer. The essential parts of a coal-gas plant are shown in Figure 35. The coal is distilled in a -shaped retorts, made of fire clay and about eight feet long. Six or more retorts are arranged in tiers form- ing a group or bench, so that all the retorts of a bench can be heated by a single fire usually of coke. Several benches placed end to end constitute a stack. The retorts are heated red hot, and about two hun- dred pounds of coal are evenly distributed on the bottom of each retort with a long iron scoop, and the mouth is quickly and tightly closed by an iron lid. The distillation continues from four to six hours, during which the temperature often reaches 1200 C. The lid is then removed, the red-hot coke is pushed or raked out, and another charge of coal is quickly introduced. The coke is quenched with water to prevent fur- ther combustion. Some of it is used for heating the retorts, but a part is sold. The volatile products pass from each retort up through a standpipe, down the dip pipe, and bubble through water into the hydraulic main. This is a horizontal, half-round pipe extending the whole length of the stack. Here some of the tar is deposited and ammonium compounds are dissolved by the water which flows constantly through the main. This water is kept at the same level and acts as a " seal " to prevent the gas from passing back into the retorts. The ammoniacal liquor and tar flow into a tar well. From the hydraulic main the gas which is hot and impure passes Illuminating Gas. 211 SJ.UOJ.3U 212 Descriptive Chemistry. into the condenser. This is a series of vertical iron pipes, several hundred feet long. They are connected at the top, but they open at the bottom into a series of boxes so constructed that the gas must pass through the entire length of the pipes, while the tar and ammoniacal liquor flow into 'the tar well. The main object of the condenser is to cool the gas slowly and condense and remove the tar. An exhauster, in most plants, draws or forces the gas from the hydraulic main through the condenser into the scrubber and onward through the purifiers into the gas holder. The exhauster also reduces the pressure in the retorts and regulates the pressure in the holder (see below) . The scrubber is a washing machine. Its purpose is to remove the remaining ammonia, part of the carbon dioxide, and hydrogen sulphide gas, and the last traces of tar. Scrubbers vary in construction. One form is a double tower filled with wooden slats or with trays covered with coke or pebbles over which ammoniacal liquor slowly trickles in the first part and pure water in the second. The gas enters at the bottom, meets the descending liquid, and is thoroughly washed. Another form widely used consists of a cylindrical vessel in which numerous wooden slats revolve in compartments and dip into am- moniacal liquor or water at the bottom. The liquid forms a film on the slats and absorbs the ammonia and other gases, while the resulting solution mixes with liquor at the bottom and flows into the proper well. Sometimes a separate tar extractor is connected with the scrubber. This is a tower filled with perforated plates, which catch and remove the tar mechanically as the gas passes through into the scrubber. From the scrubber the gas passes into the purifiers. Their FIG. 36. -Slat frame (or grid) used in chief purpose is to remove the the lime purifier. remaining carbon dioxide and sul- phur compounds. They are shal- low, rectangular iron boxes provided with slat frames loosely covered with lime (Fig. 36). In some plants iron oxide is used as the purifying material. The purified gas next passes through a large meter, which records its volume, into a gas holder. The holder is an enormous, cylindrical, iron tank in which the gas is stored. It floats in a cistern of water, and rises or falls as the gas enters or leaves. Weights and the pressure Water Gas. 213 from the exhauster so balance it that it exerts just enough pressure to force the gas through the pipes to the consumer. A ton of good coal yields about 10,000 cubic feet of gas, 1400 pounds of coke, 120 pounds of tar, 20 gallons of ammoniacal liquor, and a vary- ing amount of gas carbon. The coke is a valuable fuel and finds a ready sale. The tar, or coal tar as it is often called, collected from the hydraulic main and condenser, is a thick, black, foul-smelling liquid. It was formerly thrown away. Some is used for preserving timber, making tarred paper and concrete, and as a protective paint. Most of it is now separated by distillation into its more important constituents, especially benzene (C 6 H C ) . These carbon compounds and their numer- ous derivatives appear in commerce as oils, medicines, dyestufFs, flavors, perfumes, and other useful products. The ammoniacal liquor from the hydraulic main, condenser, and scrubber is the source of ammonia and its compounds. Gas carbon is the hard deposit which collects on the inside of the retort, and is used in the electrical industries (see Gas Carbon). The sale of these by-products reduces the cost of making the coal gas. Water Gas is made by forcing steam through a mass of red-hot coal and mixing the gaseous product with hot gases obtained from oil. The essential parts of the apparatus are shown in Figure 37. Air is forced through the coal fire in the generator, and the hot gases which are produced pass down the carburetor, up into the super- heater, and escape through its top into the open air. This operation lasts about four minutes, and is called the " blow." It heats the fire brick inside the carburetor and superheater intensely hot, air often being forced in to raise the temperature. The air valves and the top of the superheater are now closed, and the " run " begins, which lasts about six minutes. Steam is forced into the generator at the bottom. In passing through the mass of incandescent carbon the steam and carbon interact thus C + H 2 O CO + H 2 Carbon Steam Carbon Monoxide Hydrogen This mixture of hydrogen and carbon monoxide burns with a feeble flame, and before it can be used as an illuminating gas it must be 214 Descriptive Chemistry. Characteristics of Illuminating Gases. 215 enriched with gases which are illuminants. Therefore, the mixed gases pass to the top of the carburetor, where they meet a spray of oil. And as the gaseous mixture passes down the carburetor and up the super- heater, the hydrocarbons of the oil are transformed by the intense heat into hydrocarbons that do not liquefy when the gas is cooled. The ad- dition of hydrocarbons is called carbureting. From the superheater the water gas passes through the purifying apparatus into a holder. Water gas is seldom burned alone, but is usually mixed with 60 or 70 per cent of coal gas. This mixture is popu- larly called " illuminating gas." Owing to the high percen- tage of carbon monoxide, water gas and gases containing it are poisonous. Characteristics of Illuminating Gases. Both coal gas and water gas have a disagreeable odor. They are mix- tures having a composition which varies with the coal used, the temperature reached, and the degree of purifica- tion attained. The following table shows the average COMPOSITION OF ILLUMINATING GASES. CONSTITUENTS. COAL GAS. WATER GAS. Marsh 2fas . . JA C 19 8 Ethylene (and other illuminants) Hydrogen J^fO 5.0 AQ O 16.6 52 I Carbon monoxide ... 7 2 J^.l 26 I Carbon dioxide I i 30 5.2 M 2.4. Both kinds of illuminating gas may contain a little oxygen, and traces of ammonia and hydrogen sulphide gases. Nitrogen and the last portions of carbon dioxide are impurities not easily removed. Marsh gas, hydrogen, and carbon monoxide burn with a feeble (non- yellow) flame, and are often called diluents ; they furnish heat, but no light. 216 Descriptive Chemistry. The luminosity of illuminating gas depends mainly upon the presence of hydrocarbons containing a relatively large proportion of carbon. Acetylene gas, which gives such a brilliant light, consists almost wholly of this hydro- carbon containing 90 per cent of carbon. The most im- portant illuminants in coal gas and water gas are ethylene and similar hydrocarbons, acetylene, and benzene (C 6 H 6 ). The commercial value of an illuminating gas depends upon its illu- minating power. This property is measured by a photometer and is expressed in * candles." The determination is made by comparing the light produced by burning the gas in a standard burner at the rate of five cubic feet an hour with the light produced by a standard wax candle burning at the rate of 120 grains (7.77 gm.) an hour. If the gas flame is 20 times brighter than the candle flame, then the candle power of the gas is 20. The candle power of ordinary coal gas is about 17, and that of water gas is about 25. Ordinary illuminating gas has a candle power of about 20, since it is usually a mixture of coal gas and water gas. Flame. A flame is a mass of burning gas. Ordinarily it is gas combining chemically with the oxygen of the air. In the illuminating gas flame the gas itself is burning in the air. In a lamp flame the gas which burns comes from the oil which is drawn up the wick by capillary attraction, and then volatilized by the heat. Similarly, in a candle flame the burning gas comes from the melted wax. The flame produced by most burning hydrocarbons is yellowish white. The hydrocarbon flame has several distinct parts, though the structure of the flame is essentially the same, whether produced by burning illuminating gas, kerosene oil, or can- dle wax. The candle flame may be taken as the type. An examination of the enlarged vertical section shown in Fig- ure 38 reveals four somewhat conical portions, (i) Around the wick there is a black cone (A), filled with combustible Flame. 217 FIG. 38. Candle flame. gases formed from the melted wax. They do not burn be- cause no oxygen is present. With a glass tube of fine bore it is possible to draw off these gases from a large flame and light them at the upper end of the tube. (2) Around the lower part of the dark cone is a faint bluish cup-shaped part (>, B). It is the lower por- tion of the exterior cone where complete combustion of the gases occurs, since plenty of oxygen from the air reaches this portion. (3) Above the dark cone is the luminous portion (C). It is the largest and most important part of the flame. It is popu- larly spoken of as " the flame." Combus- tion is incomplete here, because little or no oxygen can pass through the exterior cone. The tempera- ture is high, however, and the hydrocarbons undergo complex changes. Acetylene is probably formed. The most characteristic change is the liberation of small par- ticles of carbon. This liberated carbon heated to incan- descence by the burning gases makes the flame luminous. The carbon glows but does not burn up, because little or no oxygen is present. A crayon or glass rod held in this part of the flame is at once coated with soot, which consists of fine particles of carbon. The exterior cone (D, D) is almost invisible. Here combustion is complete, because the oxygen of the air changes all the carbon into carbon dioxide. That this is the hottest region may be easily shown by pressing a piece of stiff white paper for an instant down upon the flame almost to the wick. The paper FIG. 39 .- Paper wi ^ fr Q charred by the outer part of the charred by a can- dle flame. flame, as shown in Figure 39. 2i 8 Descriptive Chemistry. These four portions may be found in all luminous hydrocarbon flames, whatever the shape. An ordinary gas flame is flattened by forc- ing the gas flame through a narrow slit in the burner, so that the flame will give more light. The blue part is easily seen, however, when the gas flame is turned low or looked at through a small opening ; the dark and yellow parts are always visible the latter being intentionally en- larged. The flat or circular flame of an oil lamp likewise presents the same characteristics. The gaseous products of the combustion of hydrocarbons are water vapor and carbon dioxide. A bottle in which a candle is burning has, at first, a deposit of moisture on the inside ; and if the candle is removed and limewater added, the presence of carbon dioxide is shown by the milkiness of the limewater. The oxygen needed by the burning hydro- carbons is obtained from the air. If not enough oxygen is present, the flame smokes, i.e. the carbon is thrown off into the air before the particles are heated hot enough to glow. All oil lamps are so constructed that air enters the burner below the flame. Large oil lamps have a central opening through which a large volume of air passes up inside the circular flame. Otherwise the lamp would burn with a very smoky flame. The luminosity of hydrocarbon flames is affected by other things besides the presence of glowing carbon. One of these is temperature. Gases cooled before being burned give poor light. A candle flame may be cooled enough to extinguish it. Thus, if a coil of copper wire is lowered upon a candle flame, the flame smokes, loses its yellow color, and finally goes out ; but if a coil of hot wire is used, the flame burns unchanged. Gases, as well as solids and liquids, have a kindling tem- perature, i.e. a temperature to which they must be heated before they " catch fire." This temperature differs with different substances. As we lower the temperature of gases burning with a luminous flame, their luminosity decreases, and below their kindling point they will not burn. The density of the gases in the flame and of the atmosphere itself like- wise modifies luminosity. A candle flame was found by experiment to be smaller on the top of Mont Blanc than at the base. The Bunsen Burner and its Flame. 219 O Not all flames are luminous. The hydrogen flame is almost invisible, and the flames of carbon monoxide and methane are a faint blue. These flames yield no solid particles of carbon, but only gaseous products. The most common non-luminous flame is the Bunsen flame. The Bunsen Burner and its Flame. When illuminat- ing gas is mixed with air before burning, and the mixture burned in a suitable burner, a flame is produced which is non-luminous -and very hot. The temperature of the hottest part is about 1 500 C. This flame deposits no carbon, since its products are entirely gaseous. Such a flame is called the Bunsen flame, because it is produced in a burner devised by the German chemist Bunsen. This burner is constantly used in chemical laboratories as a source of heat, and modified forms have numerous uses. One form, for example, furnishes the heat in the gas range used for cooking. The parts of an ordinary Bunsen burner are shown in Figure 40. The gas enters the base and escapes through a very small opening into the long tube, which screws down upon this opening. At the lower end of the long tube there are two holes, through which air is drawn by the gas as it rushes out of the small opening. The gas and air mix as they rise in the tube, and this mixture of air and gas burns at the top of the long tube. The size of the air holes at the bottom of the long tube may be changed by a movable ring, thus FIG. 40. Parts of a Bunsen burner. 220 Descriptive Chemistry. varying the volume of the entering air. When the holes are open, the typical colorless, hot Bunsen flame is formed. The combustion of the hydrocarbons is practically com- plete. They burn up before particles of carbon are liberated, thus making the flame non-luminous and free from soot. Apparatus heated by this flame is not black- ened. The Bunsen flame may be made momentarily luminous by shaking or blowing fine particles into the flame, such as powdered charcoal dust, finely divided metals, and sodium compounds. It was formerly believed that the non-luminous character of the Bunsen flame is solely due to the complete combustion of the carbon by the oxygen of the entering air. Recent experiments have shown, how- ever, that the result is partly due to the diluting action of the nitrogen* The gas burns at top of the tube and not inside, because the proper mixture of gas and air flows out more quickly than the flame can travel back. If the gas supply is slowly decreased, the flame becomes smaller and finally disappears with a slight explosion. This change is called "striking back." It is due to the fact that the tube contains an explo- sive mixture of air and illuminating gas, through which the flame travels faster than the mixture escapes from the tube. This explosion illus- trates in a small way what often happens when a mixture of air and illuminating gas is ignited. Sometimes the flame is not extinguished, but burns within (and sometimes without) the tube. This flame has a pale color, a disagreeable odor, and deposits soot. The Bunsen flame has many characteristic properties. Its color is bluish, and the different corres have different colors. There are really three cones: (i) the blue or greenish inner one of unburned gases ; (2) the very faint blue middle one ; (3) and the outer one, which is pale blue, and represents the blue cone in the candle flame. The middle and outer cones are not always easily distinguished ; and for all practical purposes it is convenient to divide the flame into two parts, an inner cone of unburned gases Oxidizing and Reducing Flames. 221 FIG. 41. The effects of wire gauze on a Bunsen flame. and an outer cone in which all the carbon is consumed. Combustible gases may be drawn off by a tube from the inner cone and ignited. A match laid for an instant across the top of the tube is charred only at the two points where it touches the outer cone ; and a sulphur match'- suspended by a pin across the top of an unlighted burner is not kindled when the gas is first lighted. A piece of wire gauze pressed down upon the flame shows a dark central portion surrounded by a luminous ring. The flame is beneath the gauze, although the gas passes freely through it and escapes. If the gas is extinguished and then relighted above the gauze, it will burn above but not beneath (Fig. 41). The gauze cools the gas below its kindling temperature. The miner f s safety lamp invented by Davy depends upon this last principle. It is an oil lamp surrounded by a cylinder of fine wire gauze (Fig. 42). When taken into a mine where there are explosive gases (fire damp), the flame continues to burn inside, though its size and color change. The gas often enters the lamp and burns inside, but the flame within does not ignite the gases without because the wire gauze keeps them cooled below their kindling temperature. Hence an explosion is often prevented. When miners notice changes in the lamp flame, they usually seek a safe FIG. 42. One place, form of Davy's * safety lamp. Oxidizing and Reducing Flames. The outer portiqn of the Bunsen flame is called the oxidizing flame, because here the oxygen is freely given to sub- 222 Descriptive Chemistry. -~Y- A stances. The inner portion is called the reducing flame, because here the hydrocarbons withdraw oxygen. A sketch of the general relation of these flames is shown in Figure 43. A is the most effective part of the oxidizing flame, and B of the reducing flame. At A metals are oxidized, and at B oxygen compounds are reduced. Sometimes a long tube with a small opening at one end, called a blowpipe, is used to produce these flames. A tube with a flattened top is put inside the burner tube to produce a luminous flame. The tip of the blowpipe rests in or near this flame, and if air is gently and continuously blown through the blowpipe, a long, slender flame is pro- duced, called a blowpipe flame (Fig. 44). It is like the Bunsen flame as far as its oxidizing and reducing properties are concerned. The blowpipe is used in the laboratory and by jewelers and mineral- ogists. On a large scale the blowpipe flame is used to reduce or oxidize ores and to melt refractory substances (see Compound Blowpipe) . The Bunsen flame has recently been utilized in producing the Wels- bach light. The non-luminous flame heats an inverted bag or " man- tle " of oxides of rare metals, and the mantle glows with an intense light. The candle power varies from 40 to 100. This form of burner is widely used because it produces a brilliant light. EXERCISES. 1. What are hydrocarbons ? ^ Where are they found ? Name sev- eral familiar substances containing hydrocarbons. 2. Are there many hydrocarbons ? Why ? 3. What is an homologous series of hydrocarbons ? Name four such series. FIG. 43. The oxi- dizing (A) and reduc- ing (Z?) flames. FIG. 44. Blowpipe flame, showing oxidiz- ing (A) and reducing (B) parts. Exercises. 223 4. What is methane ? What other names has it ? Where is it found ? How is it usually prepared ? State its essential properties. Why is it a dangerous gas ? Illustrate your answer by an equation. 5. What other name has the methane series ? Why ? Illustrate the following terms by the paraffin series : (a) substitution, () substi- tution product, (6-) saturated hydrocarbon. 6. What is ethylene ? How is it prepared ? Where is it found ? State its properties. Give the equation expressing the combustion of ethylene. 7. Illustrate the following terms by the ethylene series : (a) unsatu- rated hydrocarbon, (b) addition product, (c) isomerism, (d} metamer- ism, (e) isomer. 8. Review the subject of calcium carbide (see Chapter X). 9. What is acetylene ? How is it formed ? How is it prepared ? Give the equation for the reaction. Summarize the properties of acety- lene. 10. Illustrate the following terms by acetylene : (a) polymerism, (^) polymer, (V) unsaturated hydrocarbon. 1 1 . Describe the acetylene (a) flame, () burner, and (c) generator. What precautions must be observed in using acetylene as an illuminant ? 12. What is (a) choke damp, () black damp, (V) marsh gas, (//) olefiant gas ? 13. What is the formula of (a) methane, () ethylene, (c) benzene ? Why is C 2 H 2 the formula of acetylene 4 ? 14. How many volumes of oxygen are needed for the combustion of one volume of (a) methane, () ethylene, and (c) of two volumes of acetylene ? What volumes of what products are formed in each case ? What law do these relations illustrate ? 15. What is petroleum ? Where is it found ? Of what is petroleum composed ? How is it obtained from the earth ? Describe briefly the refining of petroleum. 1 6. What is kerosene ? Describe its method of preparation. Define and illustrate the \ES\bfldskingpoint. 17. State the uses of (a) gasoline, (<) lubricating oils, (c) vaseline, (tf) paraffin wax. 1 8. What is natural gas ? Where is it found ? Of what is it com- posed ? For what is it used ? 19. What is coal gas ? Describe briefly its manufacture. 20. What is coal tar ? What are its uses ? 224 Descriptive Chemistry. 21. What is ammoniacal liquor ? What is its source ? How is it obtained ? For what is it used ? 22. Review (a) coke, and (b} gas carbon (see Chapter XIV). 23. What is water gas ? Describe briefly its manufacture. What is meant by " enriching " water gas ? What is producer gas ? 24. Give the equation for the interaction of carbon and steam. How many volumes of steam are needed to produce one volume of each of the products ? 25. What is illuminating gas ? State its chief properties. What are its (a} light-giving constituents, (b} diluents, (c) impurities ? Upon what does its luminosity depend ? How is this property measured and expressed ? Give two reasons why illuminating gas is dangerous. 26. What is a flame ? Illustrate your answer. Describe the struc- ture of a candle flame. What are the chief gaseous products of combus- tion ? Why do lamps sometimes smoke ? What affects the luminosity of many flames ? 27. Describe (a) the Bunsen flame, (fr) the Bunsen burner. Why is the Bunsen flame non-luminous ? Describe and explain the " strik- ing back" of the Bunsen flame. Describe the structure of the Bunsen flame. What is the miner's safety lamp, and upon what principle is it constructed ? 28. Review oxidation and reduction. 29. What is (#) an oxidizing flame ? Describe a blowpipe and its flame. For what is it used ? 30. Describe the Welsbach light. PROBLEMS. 1. Calculate the percentage composition of (a) methane (CH 4 ), () ethylene (C 2 H 4 ), and (c} acetylene (C 2 H 2 ). 2. What weight of oxygen is needed for the complete combustion of 4 gm. of ethylene ? (Equation is C 2 H 4 + 3<3 2 = 2 CO, + 2 H 2 O.) 3. What is the simplest formula of a compound having the compo- sition H = 7.69 and C = 92.3 ? 4. Calculate the molecular formula of a compound having the vapor density 38.8 and the composition C = 92.3 and H = 7.69. CHAPTER XVI. FLUORINE - BROMINE IODINE. FLUORINE, bromine, and iodine, together with chlorine, are often grouped, and called the fialogenj. They resem- ble each other in a general way, aiuT forni analogous com- pounds which have similar properties, differing mainly in degree. Halogen means " a sea-salt producer." It is applied to this group of elements because they form salts which resemble sodium chloride (common sslt or sea salt). Chlorides, bromides, and iodides are some- times called haloid salts or halides. The Greek word for salt, hals, suggested these terms. FLUORINE. Occurrence. Fluorine is the most active of all the ele- ments, and is therefore never found free in nature. It occurs abundantly in combination with calcium as fluor spar or calcium fluoride (CaF 2 ). Other native compounds are cryolite (Na 3 AlF 6 ) and apatite (CaF 2 . 3 Ca 3 (PO 4 ) 2 ). Minute quantities of combined fluorine are found in bones and blood, in the enamel of the teeth, and in sea and some mineral waters. Fluorine is named from fluor spar, which melts easily and is used as a flux to make substances flow together (hence the derivation from the Latin fluo, I flow). The Isolation of Fluorine was accomplished in 1886 by Moissan, though many unsuccessful attempts had been previously made. He decomposed hydrofluoric acid by 225 226 Descriptive Chemistry. electricity and collected the liberated fluorine. The achievement was attended with tremendous difficulties, owing to the intense activity of fluorine and its corrosive properties. The essential parts of the apparatus used by Moissan are shown in Figure 45. The U-tube, made of an alloy of platinum and iridium, is provided with tightly fitting stoppers of fluor sr (S. S) . Through the stoppers pass the elec- trodes (E, E) of platinum iridium, held in place by screw caps (C,C). Side tubes ( T, T) allow the liberated gases (fluorine and hydrogen) to be drawn off separately through platinum delivery tubes. Perfectly dry hydrofluoric acid is put into the U-tube and dry acid potassium fluoride (HKF 2 ) is added to enable the solution to conduct the current liquid hydrofluoric acid itself being a non-conductor. The U-tube is cooled to a very low tempera- ture (23 to 50 C.), and on passing a current through the apparatus fluorine is evolved at the positive electrode and hydrogen at the other. The fluorine, freed from hydro- fluoric acid vapor, was collected by Moissan at first in a platinum tube with thin fluor spar plates closing each end, so that he could look inside and examine the gas. Later he found that pure fluorine can be collected in glass tubes, since it attacks glass only very slowly. Properties. Fluorine has a sharp odor and a greenish yellow color, but lighter and more yellowish than chlorine. Its density is 1.265 ( an " = 0- Subjected to pressure and a very low temperature, it condenses to a pale yellow liquid, which boils at 187 C. The pure gas can be liquefied in a glass vessel. Chemically, fluorine is intensely active. Hydrogen, bromine, iodine, sulphur, phosphorus, carbon, silicon, and boron take fire in it. Oxygen, nitrogen, and argon do not unite with it. Most metals burn in it, form- FlG. 45. Moissan's ap- paratus for preparing flu- orine. Fluorine Bromine Iodine. 227 ing fluorides. Gold and platinum are not attacked by it below red heat. Copper becomes coated with copper fluor- ide, which protects the metal, so that copper vessels may be used as fluorine generators. Moissan used a copper U-tube to prepare large volumes. Water is decomposed by it at ordinary temperatures, owing to the intense attrac- tion between hydrogen and fluorine ; hydrocarbons, for a similar reason, are instantly decomposed, hydrofluoric acid and carbon fluorides being the products. The exhaustive work of Moissan shows that fluorine, though more active than the other halogens, is similar to them, and should be regarded as the first member of that group. Hydrofluoric Acid, HF, is the compound of fluorine corresponding to hydrochloric acid. It is prepared by the interaction of a fluoride and concentrated sulphuric acid. Calcium fluoride is usually used, and the experi- ment is performed in a lead dish. The chemical change is represented thus CaF 2 + H 2 SO 4 = 2HF + CaSO 4 Calcium Fluoride Sulphuric Acid Hydrofluoric Acid Calcium Sulphate Hydrofluoric acid, like hydrochloric acid, is a colorless gas, which fumes in the air and dissolves in water, the solution being the commercial hydrofluoric acid. Both gas and liquid are dangerous substances. The gas is extremely poisonous, and the liquid, if dropped on the skin, produces terrible sores. Owing to its corro- sive action the acid is preserved and sold in platinum, rubber, or wax bottles. The acid and the moist gas attack glass, and are used extensively in etching. The glass is coated with wax, and the design to be etched is scratched through the wax. The glass is the*n exposed to the gas or the liquid, which attacks the exposed places. When the 228 Descriptive Chemistry. wax is removed, a permanent etching like the design is visible. Glass is an artificial compound of silicon a silicate. The corrosive action of hydrofluoric acid upon glass is due to the ease with which the acid decomposes glass and forms with the silicon a volatile compound, called silicon tetrafluoride (SiF 4 ). Since silicon dioxide (or sand) is the essential constituent of the mixture from which glass is made, the equation for etching glass may be written thus SiO 2 + 4HF = SiF 4 + 2 H 2 O Silicon Hydrofluoric Silicon Dioxide Acid Tetrafluoride Scales on thermometers and on other graduated glass instruments are etched with hydrofluoric acid. The vapor density of hydrofluoric acid gas indicates that its formula is HF at high temperature, but H 2 F 2 at lower temperatures (30 C.). BROMINE. Occurrence. Bromine is never found free in nature on account of its chemical "activity. Bromides are widely distributed, especially magnesium bromide. The salt springs of Ohio, West Virginia, Pennsylvania, and Michi- gan, and the salt deposits at Stassfurt in Germany furnish the main supply of the element. Sea water, Chili salt- peter (NaNO 3 ), and certain seaweeds contain a small quantity of combined bromine. Preparation. Bromine is obtained from its compounds by treatment with chlorine, or with sulphuric acid and manganese dioxide. In the laboratory, bromine is pre- pared by heating potassium bromide with manganese dioxide and sulphuric a'cid in a glass vessel. The bromine is easily liberated as a dense, brown vapor, which often Fluorine Bromine Iodine. 229 condenses to a liquid and runs down the walls of the vessel. The chemical change is represented thus - 2 KBr + 2 H 2 SO 4 + MnO 2 = Br 2 + MnSO 4 + K 2 SO 4 + 2 H 2 O Potassium Sulphuric Manganese Bro- Manganese Potassium Water Bromide Acid Dioxide mine Sulphate Sulphate Bromine is sometimes prepared by treating a bromide with manganese dioxide and hydrochloric acid. The source of commercial bromine in the United States is " bittern " a concentrated liquid left after salt is crystallized from brine. In the continuous process the hot bittern flows down a large tower filled with broken brick or burned clay balls ; chlorine gas and steam forced in at the bottom meet the bittern and liberate the bromine, which passes as a vapor out of the top into a condenser. The main chemical change is represented thus MgBr 2 + C1 2 - Br 2 -f MgCl 2 Magnesium Bromide Chlorine Bromine Magnesium Chloride In the periodic process, used chiefly in the United States, a huge stone still is charged with manganese dioxide, hot bittern, and sulphuric acid, and heated by steam. The bromine distills into a condenser, as in the other process. Sometimes potassium chlorate is used as the oxidizing agent. Properties.VtJ^romineis a heavy, reddish brown liquid at the ordinary T^njgCIamel Its specific gravity is about three. It is a volatile liqu!?f, boiling at about 59 C. The vapor, which is given off freely, has a disagreeable, suffo- cating odor. This property suggested the name bromine (from the Greek word bromos, a stench). It is poisonous, and burns the flesh frightfully. Bromine is somewhat soluble in water. The solution, called bromine water, has a brown color, and when cooled deposits a crystalline hydrate (Br 2 . 10 H 2 O). Many other properties of bromine are similar to those of chlorine. Thus, it combines with metals and other elements ; it also bleaches. 230 Descriptive Chemistry. Compounds of Bromine are similar to those of chlorine. Hydrobro- mic acid (HBr) is a colorless, pungent gas, which fumes in the air and dissolves freely in water, forming the solution usually called hydrobromic ' acid. Its other properties closely resemble those of hydrochloric acid. Bromides are salts of hydrobromic acid, though many are formed by direct combination with bromine. Like the chlorides, most bromides dissolve in water. Potassium bromide (KBr) is a white solid, made by decomposing iron bromide with potassium carbonate. It is used exten- sively as a medicine and in photography (in preparing silver bromide plates and films). Bromides of sodium, ammonium, and cadmium have a limited use. Miscellaneous. Bromine itself is used to make potassium bromide and other compounds, especially a class of coal tar dyes used to color pink string and to make red ink. Annually over 500,000 pounds of bromine are prepared in the United States, while Germany exports about 400,000 pounds of bromine, and 500,000 pounds of bromine compounds. Balard discovered bromine in 1826 in the mother liquor (or bittern) from brine. Liebig supposed it was chloride of iodine, and thus failed to discover it, because, as he said, he yielded to " explanations not founded on experiment." IODINE. Occurrence. Free iodine is never found in nature, but like chlorine and bromine it is combined with metals, especially sodium, potassium, or magnesium. It is widely distributed, though the quantity in any one place is small. Tobacco, water cress, cod-liver oil, oysters, and sponges con- tain minute quantities. Native iodides of silver and of mer- cury are found. The ash of some seaweeds contains from 0.5 to 1.5 per cent of its weight of iodides of sodium and potassium. Sodium iodate occurs in the deposits of salt- peter in Chili, and is now the main source of the element. Preparation. Iodine is prepared in the laboratory by a method similar to that used for bromine. Potassium iodide, manganese dioxide, and sulphuric acid are heated in a glass vessel, and the iodine appears as a violet vapor, which con- A . Fluorine Bromine Iodine. 231 denses on the upper part of the vessel into dark grayish crystals. On a commercial scale iodine is prepared from the ash of seaweeds and from the mother liquors of Chili saltpeter, (i) Along the coasts of France, Scotland, and Norway seaweed is collected and burned, usually in closed vessels. The ash is called kelp or varec. The solu- ble portions are removed by agitation with water. The 'filtered liquid is further purified, and from the final mother liquor in which the iodides are dissolved, the iodine is extracted by heating with sulphuric acid and manganese dioxide. Sometimes chlorine is used to extract the iodine. In either case the mother liquor and its added ingredients are distilled FiG. 46. Apparatus for purifying iodine. gently in an iron pot with a lead cover, which is connected with two rows of bottle-shaped condensers (Fig. 46). The iodine, which col- lects in these condensers, is purified by washing and resubliming. (2) In another process the mother liquor from the Chili saltpeter is mixed with acid sodium sulphite (HNaSO 3 ), and the precipitated iodine is collected on coarse cloth, washed, dried, and then resublimed, as described above. Courtois, a French chemist, discovered iodine, in 1812, in an attempt to prepare potassium nitrate from seaweed. Davy and Gay-Lussac established its elementary nature and discovered many of its properties. The present name was given by Davy. Properties. Iodine is a dark grayish crystalline solid, resembling graphite in luster. It crystallizes in plates which have the specific gravity 4.95. It is volatile at the 232 Descriptive Chemistry. ordinary temperature, and when gently heated the vapor which is formed has a beautiful violet color. This color suggested the name iodine (from the Greek word iodes, violetlike). The vapor is nearly nine times heavier than air, and has an odor resembling dilute chlorine, though less irritating. When the vapor is heated, its color changes from violet to deep blue, and the density decreases. Ex- periment indicates that at about 700 C. the molecules con- tain only two atoms, and as the temperature rises the molecules dissociate, until at a very high temperature the vapor consists entirely of atoms. Iodine stains the skin yellow, and turns cold starch solution blue. The presence of a minute trace of iodine may be thus detected, one part of iodine in over 400,000 parts of water producing the blue color. The exact nature of this blue compound is un- known. The presence of starch in many vegetable sub- stances can be shown by this delicate test. Iodine dissolves slightly in water, and freely in alcohol, chloroform, carbon disulphide, ether, and potassium iodide solution. The chloroform and carbon disulphide solutions are violet, but the others are brown, or even black. The chemical proper- ties of iodine resemble those of chlorine and bromine, but it is less active. Bromine and chlorine displace iodine from its compounds, chlorine and chlorine water being often used for this purpose. It combines directly with other elements and replaces some. Phosphorus bursts into a flame when mixed with iodine. Compounds of Iodine resemble the corresponding ones of chlorine and bromine. Hydriodic acid is much like hydrobromic and hydro- chloric acid, though unlike them in being a reducing agent. Iodides are salts of hydriodic acid, and like many salts they are prepared in various ways. In general behavior they are similar to bromides and chlorides. Potassium iodide (KI) is made and used like potassium bromide. lodates and periodates are known. Fluorine Bromine-^- Iodine. 233 Miscellaneous. Iodine dissolved in alcohol or in potassium iodide solution is used as an application for the skin to prevent the spread of eruptions or to reduce swellings. Iodine is used to make medicinal preparations, especially iodoform (CHI 3 ), which is used as a dressing for wounds. Large quantities of iodine are used in making aniline dyes. Potassium iodide is made in large quantities, Germany alone exporting about 150 tons of it annually. Chili annually exports over 300 tons and Norway over 160 tons of iodine and iodides. EXERCISES. 1. What elements constitute the halogen group ? Why are they so .called ? 2. How does fluorine occur in nature ? Describe briefly the isola- tion of fluorine. When was it first performed? Summarize the chief properties of fluorine. 3. How is hydrofluoric acid prepared? Give the equation for the reaction. What are its characteristic properties ? For what is it used ? 4. How is glass etched? State the essential changes. 5. What is the formula of hydrofluoric acid ? 6. How does bromine occur in nature ? What are the sources of commercial bromine ? What general method is used to prepare this element ? Describe briefly the commercial methods. State the chief properties. For what is it used ? How does this element differ from all others previously studied ? 7. Name several compounds of bromine. What is potassium bromide ? 8. Give a brief account of the discovery of (a) bromine and (b) iodine. 9. Discuss the occurrence of iodine in nature. How is iodine pre- pared (a} in the laboratory and (b) on a large sqale ? Summarize the properties of iodine. Describe the test for iodine. 10. Name several compounds of iodine. Describe potassium iodide, n. Compare hydrochloric, hydrobromic, and hydriodic acids. 12. What is the symbol of (a} fluorine, (b} chlorine, (c) bromine, (d) iodine ? What is the derivation of the name of each element ? 13. Compare the physical properties of fluorine, chlorine, bromine, and iodine. 14. What is " drug-store iodine " ? 234 Descriptive Chemistry. PROBLEMS. 1 . What is the percentage composition of (a) fluor spar (CaF 2 ) and () cryolite (Na 3 AlF 6 )? 2. How much (a) calcium sulphate and () hydrofluoric acid are formed by heating 100 gm. of fluor spar with sulphuric acid ? 3. Calculate the percentage composition of (a) potassium bromide (KBr), () potassium iodide (KI), (c} silver bromide (AgBr), and (/ Phosphine Hydriodic Acid Phosphonium Iodide 270 Descriptive Chemistry. Phosphorus Trichloride (PCI.,) is a disagreeable smelling liquid, made by the combustion of dry chlorine and phosphorus ; and phosphorus pentachloride (PC1-) is a greenish solid made by passing chlorine into a vessel containing the trichloride. Matches. Phosphorus is chiefly used in the manufac- ture of matches. Soft wood is cut by machinery into the desired shape. The cards or sticks are fixed in a frame, and one end is first dipped into melted sulphur or paraffin and then into the phosphorus mixture. The latter consists usually of different proportions of phosphorus, manganese dioxide, glue, and a little coloring matter. Manganese di- oxide may be replaced by other oxidizing agents. These matches are the ordinary friction or sulphur kind. By rubbing them on a rough surface enough heat is gener- ated to cause the phosphorus to unite with the oxygen of the oxidizing agent, and the heat thereby produced sets fire to the sulphur or paraffin, and this in turn kindles the wood. Since these matches are poisonous, and liable to take fire, their manufacture has been prohibited in some countries (e.g. Switzerland and the Netherlands). Safety matches, which replace them, contain no yellow phos- phorus. The head of this kind is usually a colored mix- ture of antimony sulphide, potassium chlorate, and glue; while the surface upon which the match must be rubbed to light is coated with a mixture of red phosphorus, glue, and powdered glass. Matches are made by machinery, several million being produced in one day. Relation of Phosphorus to Life. Phosphorus is essen- tial to the growth of plants and animals. Plants take phosphates from the soil and store up the phosphorus compounds, especially in their fruits and seeds. Animals eat this vegetable matter, assimilate the phosphorus com- pounds, and deposit them in the bones, brain, and nerve Phosphorus, Arsenic, Antimony, Bismuth. 271 tissue. Bones contain about 60 per cent of calcium phos- phate. Part of the combined phosphorus consumed by animals is rejected by them, and often finds its way back into the soil. The constant removal of phosphates by plants would soon exhaust the soil. Hence phosphorus is restored to the soil in the form of natu- ral or artificial fertilizers. Natural fertilizers are (i) stable refuse, which always contains some of the phosphates from the food originally fed to the animals ; (2) guano, which is the dried excrement and carcasses of the sea birds that once lived in vast numbers in Peru and Chili ; and (3) phosphate slag, which is a phosphorus by-product obtained in manu- facturing steel. These and bones are ground and spread upon the soil. Artificial fertilizers are made from phosphate rock. This occurs in large beds in South Carolina, Tennessee, and Florida, which yield about a million tons a year. It consists of the hardened remains of land and marine animals, and is mainly tricalcium phosphate (Ca 3 (PO 4 ) 9 ). It is insoluble in water, and must be changed into the soluble monocalcium salt (H 4 Ca(PO 4 ) 2 , so that it can be evenly distributed through the soil and easily taken up by plants. This soluble salt is called " superphos- phate of lime." When phosphate rock is treated with sulphuric acid, the changes involved may be written thus Ca 3 (PO 4 ) 2 + '2H 2 SO 4 = H 4 Ca(PO 4 ) 2 + 2CaSO 4 Tricalcium " Superphosphate Calcium Phosphate of Lime " Sulphate Ca 3 (P0 4 ) 2 + 3H 2 S0 4 = 2H 3 P0 4 + 3 CaSO 4 Phosphoric Acid Ca 3 (PO 4 ) 2 + H 2 SO 4 = H 2 Ca 2 (PO 4 ) 2 + CaSO 4 Dicalcium Phosphate The aim is to convert the crude phosphate rock into "superphos- phate," but the three reactions usually occur. The product is ground, dried, and packed in bags for the market. On standing, it may undergo " reversion," i.e. the " superphosphate " and phosphoric acid may form insoluble phosphates, thus making the fertilizer less valuable. Some- times " superphosphate " is mixed with compounds of nitrogen and of potash to produce a complete fertilizer. 272 Descriptive Chemistry. ARSENIC. Occurrence. Arsenic is found free in nature, but it usually occurs combined with sulphur or a metal, or with both. The common arsenic ores are realgar (As 2 S 2 ), orpiment (As 2 S 3 ), arsenic pyrites or mispickel (FeSAs). Arsenic trioxide or arsenolite(As 2 O 3 ) is also found. Small quantities of arsenic occur in many ores. The United States annually imports over 6,000,000 pounds of arsenic and its compounds, mainly from England and Germany. Arsenic is prepared in the laboratory by heating a mixture of arse- nious oxide and charcoal in a glass tube. The change is represented thus 2As 3 O 8 -f 6C As 4 + 6 CO Arsenious Oxide Carbon Arsenic Carbon Monoxide On a large scale it is extracted from its ores either by the above method or by roasting arsenic pyrites (FeSAs) in the absence of oxygen. Arsenic has marked properties. It is a brittle, steel-gray solid. A freshly broken piece has a metallic luster, which disappears slowly in a moist atmosphere. It tends to crystallize. The specific gravity is from 5.62 to 5.96. Heated in the air, it volatilizes without melting, and the vapor has an odor like garlic. At about 180 C. it burns in the air with a bluish flame, forming the white oxide (As 2 O 3 ). Arsenic molecules, like those of phosphorus, contain four atoms. In some respects arsenic resembles both metals and non-metals. It is used to harden the lead which is made into shot. Arsenious Oxide or Arsenic Trioxide, As 2 O 3 , is the most important compound of arsenic, and is often called simply " arsenic" or "white arsenic." It is found free in nature, but is usually manufactured by roasting arsenic ores. There are two common varieties, a white, granular powder and an amorphous, glasslike solid. It has no odor, a faint, metallic taste, dissolves slightly in cold water, but readily in hot hydrochloric acid. Arsenic trioxide is a Phosphorus, Arsenic, Antimony, Bismuth. 273 rank poison. The antidote is fresh ferric hydroxide, which is made by adding ammonium hydroxide to a ferric salt, e.g. ferric chloride. Small doses (2 to 3 grains) are usually fatal, but by habitual use the system appropriates larger doses without ill effects. Workmen in arsenic factories often accidentally swallow with impunity quantities which would ordinarily prove fatal. It is used for making pig- ments for green paints, for fly and rat poison, in mak- ing glass, arsenic compounds, in calico printing, and in preserving skins. As a medicine it is used to purify the blood. Other Arsenic Compounds. The native mineral orpiment (As 2 S 3 ) is used in making a- yellow paint, and realgar (As 2 S 2 ) a red paint. Scheele's green is chiefly copper arsenite (HCuAsO 3 ), and was formerly used to make a cheap green paint and to color wall paper. The com- plex arsenic compound Paris green is a light green powder ; owing to its poisonous character it is used to exterminate potato bugs and other insects. Arsenic forms acids analogous to the acids of phosphorus, though they are less important. The salts sodium arsenate (HNa 2 AsO 4 ) and arsenite (NaAsO 2 ) are used in dyeing. The formation of the yel- low sulphide (As 2 S 3 ) by passing hydrogen sulphide into an arsenic solution containing hydrochloric acid is the usual test for arsenic. Marsh's Test for Arsenic. Arsenic itself is not poisonous, but its compounds are among the most poisonous substances known. For- tunately, combined arsenic is easily detected by a simple method, called Marsh's test. An apparatus for generating hydrogen is provided with a hard glass horizontal delivery tube, narrowed in places and drawn to a point. Pure zinc, pure dilute sulphuric acid, and the arsenic solution are put in the generator. Hydrogen and gaseous hydrogen arsenide (or arsine (AsHo) ) are formed. If this mixture is lighted at the end of the delivery tube, metallic arsenic is deposited as a black coating on cold porcelain held in the flame ; or if the tube is heated in front of a narrow place, arsenic is deposited at this point. This deposit dissolves in sodium hypochlorite solution, but a deposit of antimony, similarly produced, does not dissolve. By this delicate test the merest trace of arsenic is readily and positively detected. 274 Descriptive Chemistry. ANTIMONY. Occurrence of Antimony. Small quantities of free anti- mony are found. The most common ore is stibnite (Sb 2 S 3 ), which occurs in Japan, Austria-Hungary, France, Algeria, Italy, Mexico, and Turkey. Large deposits in California and Nevada are now utilized, about 3,000,000 pounds being annually produced. Stibnite was known in the fifteenth century. The Latin name of antimony is stibium, from stibnite, which gives the symbol of the element, Sb. Antimony is prepared on a large scale by two methods. In one the sulphide is roasted, and the oxide thus formed is reduced with charcoal. Equations representing the main changes are 2Sb 2 S 3 + 9O 2 = 2SbO 3 4 6 SO 2 Antimony Sulphide Oxygen Antimony Oxide Sulphur Dioxide 2Sb 2 O 3 + 3C 4Sb + 3CO 2 The other method consists in heating the sulphide with iron, the equation for the chemical change being Sb 2 S 3 + 3Fe = 2Sb -f 3 FeS Antimony Sulphide Iron Antimony Iron Sulphide Antimony has interesting properties. It is a silver white, crystal- line, brittle solid. Its specific gravity is 6.7. At ordinary temperatures antimony does not tarnish in the air, but when heated, it burns with a bluish flame, forming the white, powdery antimony trioxide (Sb 2 O.,). Powdered antimony burns brilliantly when added to chlorine, bromine, or iodine. Nitric acid oxidizes it, and aqua regia dissolves it. Anti- mony melts at about 450 C. It expands on cooling, and is therefore one constituent of type metal (see Alloys of Lead). Compounds of Antimony. Antimony forms stibine (SbH 3 ), which is analogous to ammonia (NH 3 ) and arsine (AsH 3 ), pyro- and meta- acids, the oxides, Sb 2 O 3 and Sb 2 O-, and halogen compounds. It also forms complex compounds in which antimony acts as a metal. Tartar emetic is potassium antimonyl tartrate (KSbO .C 4 H 4 O 6 ). It is used as a medicine and as a mordant in dyeing cotton. Antimony trisulphide Phosphorus, Arsenic, Antimony, Bismuth. 275 (Sb.,S 3 ) is a reddish solid, formed by passing hydrogen sulphide gas into a solution of antimony the test for antimony. The sulphide is used in making the red rubber tubing and stoppers used in the labora- tory. The chlorides (SbCl 3 and SbCl 3 ) are formed by the action of chlorine upon the metal ; with water they form the white solids called oxychlorides, e.g. SbOCl. The formation of antimony oxychloride is sometimes used as a test for antimony, but the more common test is the formation of the reddish orange sulphide (Sb 2 S 3 ). BISMUTH. Bismuth is usually found in the native state, though it is not abundant nor widely distributed. The oxide (Bi 2 O 3 ), or bismite, the carbonate ((BiO) 2 CO 3 .H 2 O), or bismutite, and the sulphide (Bi 2 S 3 ), or bismuthinite, are the common ores. The world's supply comes from Saxony. Bismuth is prepared from the native metal by melting it on an inclined plate and allowing it to drain away from the solid impurities. Sometimes the sulphide is roasted, and the resulting oxide is reduced with charcoal, as in the case of antimony. Bismuth has characteristic properties. It is a grayish white metal with a reddish tinge. Like antimony, it is very brittle. It does not tarnish in dry air, but it grows dull in moist air ; and when heated in air it burns with a bluish flame, forming the yellowish oxide (Bi 2 O 3 ). Its specific gravity is about 9.9. Hydrochloric acid does not readily attack it, but nitric acid converts it into a nitrate, and hot sulphuric acid into a sulphate. Bismuth melts at about 270 C. But a mixture of bismuth, lead, and tin melts at a low temperature. For example, Newton's metal melts at 95 C. and Rose's metal at 100 C. ; while Wood's metal, which con- tains cadmium, melts at only 66 C.-yi C. These metallic mixtures are called fusible metals. They are used in making casts of wood cuts; but more often (i) as safety plugs in steam boilers to prevent explosions, (2) as a fuse in electrical apparatus to prevent a short cir- cuit, and (3) to hold in place fireproof doors and the valves in the automatic sprinkling apparatus now placed in large buildings. Compounds of Bismuth. Bismuth forms no compound with hydro- gen. There are three oxides. Bismuth trioxide (Bi 2 O 3 ) is yellowish, 2j6 Descriptive Chemistry. the pentoxide (Bi 2 O 5 ) is orange red, and the dioxide (Bi 2 (X) is black. Bismuth trioxide is used to fix the gilding on porcelain. The trichloride (BiCl t3 ) is formed by the action of chlorine upon bismuth, or by treat- ing bismuth with aqua regia. With an excess of water the trichloride forms the oxychloride (BiOCl), which is a pearl-white powder, insoluble in water. The formation of the oxychloride is the usual test for bis- muth. Bismuth, being a metal, forms hydroxides (Bi(OH) 3 and BiO.OH). Normal bismuth nitrate (Bi(NO. ? ) ;i ), treated with hot water, forms basic bismuth nitrate (Bi(OH) 2 NCX or BiONO 3 ). The latter, often called subnitrate of bismuth, is a white powder used as a medicine for dyspepsia and as a cosmetic. EXERCISES. 1 . What is the symbol and atomic weight of phosphorus ? Give a brief history of this element. Why is it so named ? 2. Discuss the occurrence of phosphorus. 3. Describe the manufacture of phosphorus (a) from a .phosphate and sulphuric acid, and (b) by the electric method. How is it purified ? 4. Summarize the properties of (a) ordinary phosphorus, and () red phosphorus. 5. Describe briefly (a) the oxides of phosphorus, (b} orthophos- phoric acid, (c) metaphosphoric acid, (d} pyrophosphoric acid, (e) phos- phine, (/) the phosphorus chlorides. 6. What is (a) tricalcium phosphate, (b} microcosmic salt, (c} " acid phosphate " ? 7. Describe the manufacture of (a) sulphur matches, and (b} safety matches. 8. Discuss the relation of phosphorus to life. 9. What is a fertilizer ? Name three natural fertilizers. Describe the manufacture of artificial fertilizer. What is a complete fertilizer ? 10. What is the symbol and atomic weight of arsenic ? 1 1 . Name several ores of arsenic. With what metals is arsenic often associated ? 12. Describe the preparation and state the properties of the arsenic. 13. What is the formula of arsenic trioxide ? By what other names is it known ? Summarize its properties. For what is it used ? What is the antidote for arsenic poisoning ? Phosphorus, Arsenic, Antimony, Bismuth. 277 14. What is (a) Paris green, (6) orpiment, (V) realgar ? For what is each used ? 15. Describe Marsh's test for arsenic. 1 6. What is the symbol and atomic weight of antimony ? 17. In what forms does antimony occur and where is it found ? De- scribe its preparation. State its chief properties. 18. What is tartar emetic ? For what is it used ? 19. Describe the test for antimony. 20. What is the symbol of bismuth ? How does it occur and where is it found ? Describe its preparation. State its properties. 21. State the relation of bismuth hydroxide to bismuth subnitrate. Describe the latter. PROBLEMS. 1. Calculate the percentage composition of (a) sodium phosphate (Na 3 PO 4 ), () dihydrogen phosphate (H 2 NaPO 4 ), (V) disodium phos- phate (HNa 2 PO 4 ), O/) microcosmic salt (HNaNH 4 PO 4 ). 2. How much phosphorus is needed to remove the oxygen from a liter of air ? (Assume (i)2P + 5O = P 2 O 5 and (2) air is 20 per cent oxygen.) 3. How much phosphorus is there in a ton (2000 Ib.) of bone ash (Ca 3 (P0 4 ) 2 )? 4. If a skeleton weighs 25 Ib. and contains 60 per cent calcium phosphate, how much phosphorus does it contain ? 5. What is the weight of a cylindrical stick of ordinary phosphorus 10 cm. long and 15 mm. in diameter ? (Suggestion. What is the spe- cific gravity of phosphorus ?) 6. Calculate the percentage composition of (a) orpiment (As 3 S 3 ), ($) realgar (As S 2 ), (^) white arsenic (As 2 O 3 ). 7. What is the weight of a piece of antimony 25 cm. long, 15 cm. wide, and 2 mm. thick ? CHAPTER XX. METALS. Introduction. The elements studied thus far are chiefly non-metals. Metals, however, have been mentioned, and many of their properties have been discussed. It is the purpose of the present chapter to review these properties and prepare the way for a fuller treatment of the metals. Metals and Non-metals. Many years ago the chem- ical elements were divided into two classes, called metals and non-metals. The division was based largely on the physical properties of the elements. The opaque, lustrous, more or less heavy, hard, ductile, malleable, tenacious solids were called metals. All gases and the solids such as carbon, sulphur, phosphorus, and iodine were called non-metals. No such sharp dividing line, however, can be drawn between metals and non-metals. Some, of course, have pronounced properties, like the non-metal sulphur and the metal iron. These are typical. But a few have variable properties. Sometimes they act as metals and at other times as non-metals. Antimony and arsenic belong to this border-line class ; they are sometimes called the metalloids. The classification into metals and non-metals is no longer accurate, but it is very convenient. The use in common life of the words metallic and metal seldom leads to confusion. Properties of Metals. The physical properties of metals are familiar, though variable between wide limits. 278 Metals. 279 All have a metallic luster, i.e. the marked property of reflecting light from their polished or untarnished surfaces. All are opaque except very thin films of gold. The color of many is white, though the tint varies. Thus silver, sodium, aluminium, mercury, magnesium, iron, and tin are nearly pure white, and bismuth is reddish white. Copper is the only red metal, and gold the only yellow one, which is an element. Most metals are malleable and ductile, i.e. they may be hammered or rolled into sheets and drawn into wire. Gold, copper, silver, iron, platinum, and aluminium possess both these properties to a marked degree ; while lead, zinc, and tin are very malleable though not so ductile. Antimony and bismuth are brittle. The hardness of metals varies. At the ordinary temperature mercury is a liquid, sodium and lead can be cut easily with a knife, and so on through the list up to iridium, which is as hard as steel. In specific gravity, which was once thought must very high, the metals range between lithium, which has the specific gravity 0.585, and osmium, which has the specific gravity 22.48. Sodium and potassium also are lighter than water, while magnesium has the specific grav- ity 1.75, and aluminium 2.58. Metals are good conductors of heat and electricity. They also vary in this property. Silver, copper, and aluminium are the best conductors, and have therefore many practical applications. Bismuth is the poorest conductor. The distinctive property of metals is not physical, but chemical. Metals form oxides which combine with water to produce bases. Metals are the characteristic elements of bases. On the other hand, non-metals form acid-pro- ducing compounds. Occurrence of Metals. Only a few metals are found free in the earth's crust, and these are seldom pure. Of 280 Descriptive Chemistry. the six metals known to the ancients, gold, copper, silver, tin, iron, and lead, only gold and copper are found free. The solid elements and their compounds which occur in the earth's crust are called minerals. And those minerals from which metals can be profitably extracted are called ores. The most abundant classes of ores are oxides, sul- phides, carbonates, and hydroxides. Lead, zinc, mercury, and silver sulphides are abundant. Besides native copper, the sulphide and carbonate are found. Iron occurs as oxide, carbonate, hydroxide, and sulphide. Many ores contain arsenic. Some ores are very complex. Preparation of Metals. The series of operations by which useful metals are extracted from their ores is called metallurgy. It includes preliminary treatment, smelting, electrolysis, refining, and other operations necessary to change the ore into a metal ready for manufacture into useful articles. The object of the preliminary treat- ment is to prepare the ore for smelting or for a similar operation by which the metal is obtained in a state adapted for further purification or refining. The ore as it comes from the mine is usually mixed with earthy matter or rock called gangue. This impurity is removed by me- chanical or chemical processes, sometimes by both. The mechanical process illustrates one kind of preliminary treat- ment. The ore is first crushed in a stamp mill. This is a huge, heavy mortar and pestle. The pestle or stamp falls repeatedly upon the ore, which is slowly fed into the mortar or die. A current of water (or air) forces the fine particles out of the mortar through a sieve. The lighter particles of the impurities are washed away, and the metallic grains are extracted by another mechanical operation, though chemical processes are frequently employed, especially with inferior ores. This separation of the valuable part Metals. 281 of the ore from the gangue, and reducing it to a smaller bulk is often called ore dressing or concentration. Copper is extracted from the Lake Superior ores mainly by this method of preliminary treatment. Gold and silver ores are treated this way, and then ex- tracted from the slime by mercury. The latter operation is called amalgamation. The most common method of extracting metals from their ores is by smelting. The process varies with the kind and composition of the ore. Essentially, it consists in heating a mixture of the ore and coke (or coal) in a furnace. The ores used must, as a rule, be oxides. Sulphides, hydroxides, and carbonates are first roasted or calcined to convert them into oxides. The essential chemical change in smelting is a reduction of the oxide by the carbon. The carbon and oxygen unite and pass off as a gas, leaving the metal to run out at the bot- tom. Limestone, or a similar substance, called a flux, is added to the mixture, if necessary, to facilitate the melting and to assist in removing the impurities as a glassy sub- stance, called slag. The operation is conducted in differ- ent kinds of furnaces. Iron, for example, is smelted in a huge upright furnace called a blast furnace (Fig. 72), because a current of air is forced through the melted mass to facilitate the fusion and chemical changes. In such a furnace the fuel and ore are in direct contact. When this is undesirable, the reverberatory furnace is used (Fig. 54). As the figure shows, in this furnace the flame is reflected or reverberated upon the ore under treatment. In this kind of furnace the ore may be oxidized or reduced with- out coming in contact with the fuel. Some ores demand special methods, which will be described in connection with these metals. Electrolysis is used to extract some metals, especially aluminium. Other metals, notably 282 Descriptive Chemistry. FIG. 54. Reverberatory furnace. Tue tire copper, are purified by electrolysis. A few met- als are extracted by a wet process. That is, the ores are dissolved, and the metal is precipi- tated by adding some substance or by elec- i ivj. 5*j- - AV<- v^i u\,i aiwi j mi lit L\^... i nt; me -I . T*!. * X" " burns on the grate, G, and the long flame trolySlS. 1 huS, interior which passes over the bridge, E, is reflected Pjold Ores are dissolved down by the sloping roof upon the contents . , of the furnace. Gases escape through /. The by treatment With potas- charge, which rests upon B, does not come sium CV^nide and the in contact with the fuel, but is oxidized or . . reduced by the flame. gold IS then precipitated by zinc. Alloys are mixtures or compounds of two or more metals. Some fused metals mix in all proportions, while others seem to form definite compounds. The properties of alloys vary with the constituents and their properties. Some alloys, especially those of copper and of lead, have many industrial uses. Alloys in which mercury is a con- stituent are called amalgams. EXERCISES. 1. Define the terms metal and non-metal as they are ordinarily used. Name six or more examples of each class. Define and illustrate the term metalloid. Why is this classification inaccurate? 2. State the familiar physical properties of metals. Define (a) metallic luster, (b} malleable, (c) ductile, (d) specific gravity. 3. How does the color of metals differ from their luster? Name five metals which are white. What color has (a) gold, (6) copper, (c) zinc, (a) lead, (e) iron? 4. What metals are brittle? Malleable? Soft? Hard? Heavy? Light? What metals conduct electricity well? 5. What is the distinctive chemical property of metals? Of nor>' metals? Illustrate your answer. Metals. 283 6. What metals are often found free in nature? Define and illus- trate the terms (a) mineral, and (b) ore. What are the most abundant classes of ores? 7. What metals occur abundantly as (a) sulphides, (b) oxides, (V) carbonates ? 8. Define metallurgy. W T hat general operations does it include? 9. What is the object of the preliminary treatment of ores? How is it accomplished mechanically? Define (a) gangue, (<) concentra- tion, (c) amalgamation. What metal is often extracted (a) mechanic- ally, (b) by amalgamation ? 10. Define smelting. What fundamental chemical change does it usually involve? Define and illustrate (a) calcination, ($) flux, (c) slag. n. Describe (a} a reverberatory furnace, and (b) a blast furnace. What is their essential difference ? 12. What is the wet process of extracting ores? 13. What are (a) alloys, (b) amalgams? PROBLEMS. 1. What is the specific gravity of gold, if a piece weighs 4.676 gm. in air,' and loses 0.244 gn. when weighed in water? (Note. Specific gravity equals the weight in air divided by the loss of weight in water.) 2. A piece of aluminium weighs 150 gm. in air and 75 gm. in water. What is its specific gravity? 3. A piece of iron weighs 292.8 gm. in air and 255.3 gm. in water. What is its specific gravity? 4. A piece of copper weighing 50 gm. in air lost 5.6 gm. when weighed in water. What is its specific gravity? 5. A piece of lead pipe weighs 158.9 gm. in air and 144.9 & m - ^ n water. Calculate the specific gravity. CHAPTER XXL SODIUM, POTASSIUM, AND LITHIUM. Introduction. Sodium and potassium, and the rare elements lithium, rubidium, and caesium, form a natural group, known as the alkali metals. The different elements and their corresponding compounds resemble each other closely. Sodium and potassium were discovered by Sir Humphry Davy in 1807 by the electrolysis of their hydroxides. Bunsen, by means of the spectroscope, discovered lithium in 1855, caesium in 1860, and rubidium in 1861. SODIUM. Occurrence. Sodium is not found free. Sodium chlo- ride and sodium nitrate are the most abundant compounds. Many rocks, plants, and mineral waters contain combined sodium. About 2.5 per cent of the earth's crust is sodium. The symbol of sodium, Na, is from the Latin word natrium, which in turn comes from the Greek word natron, an old name of sodium carbonate. Preparation. Sodium is now manufactured on a large scale by the electrolysis of fused sodium hydroxide. This method was used by Davy in 1807 to isolate sodium, but its commercial success was only recently made possible by Castner. Figure 55 is a sketch of the apparatus used. The body of the steel cylinder, S, rests within a heated flue. Hence the sodium hydroxide is solid in the neck, B, and serves to protect the joint made by the iron cathode, 284 SIR HUMPHREY DAVY 1778-1829 THE FAMOUS ENGLISH CHEViST WHOSE BRILLIANT DISCOVERIES HAVE NEVER BEEN SURPASSED Sodium, Potassium, and Lithium. 285 C, and the crucible. A, A is the iron anode. A collecting pot, P, dips into the molten caustic soda. As the electrolysis pro- ceeds, the sodium formed at C collects in P, and a wire gauze, G, G, keeps it from mix- ing with the caustic soda. The sodium is ladled out at intervals from P. The hy- drogen, which is liberated, accumulates also in P and prevents the sodium from oxi- dizing. The hydrogen some- times escapes and explodes. FIG. 55. Apparatus for the manu- facture of sodium by the electrolysis Sodium was formerly manufactured O f sodium hydroxide, by two methods, (i) Sodium car- bonate and carbon heated to a high temperature change thus Na 2 C0 3 Sodium Carbonate 2C = 2Na Carbon Sodium + 3 CO Carbon Monoxide The mixture was heated in iron retorts, and the sodium vapor, in pass- ing through a flat iron receiver, condensed to a liquid, which was col- lected under paraffin or mineral oil. (2) The other chemical method, devised by Castner in 1886, consisted essentially in heating sodium hydroxide with a mixture of iron and carbon. Probably iron carbide was the essential reducing agent, and the change might be represented thus 6NaOH + FeC 2 = 2 Na + Fe + 2 Na,CO, + 3 H 2 Sodium Hy- Iron Car- Sodium Iron Sodium Car- Hydrogen droxide bide bonate Properties. Sodium is a silver-white metal. It is so soft that it may be easily molded with the fingers and cut with a knife. It floats upon water, since its specific 286 Descriptive Chemistry. gravity is only 0.98. Heated in the air, it melts at 96 C, and at a higher temperature it burns with a brilliant yellow flame, forming the oxides Na 2 O and Na 2 O 2 . This intense yellow color is characteristic of sodium and is the usual test for the element (free or combined). In moist air the bright surface quickly tarnishes, and sodium as usually seen has a brownish coating. It is, therefore, kept under kerosene or a liquid free from water. It decomposes water at ordinary temperatures, liberating hydrogen and forming sodium hydroxide, thus Na + H 2 O - NaOH + H Sodium Water Sodium Hydroxide Hydrogen If held in one place upon water by filter paper, enough heat is generated to set fire to the hydrogen, which burns with a yellow flame, owing to the presence of volatilized sodium (see Interaction of Sodium and Water, Chapter V). If melted sodium is put into chlorine, the two elements combine with a brilliant flame, forming sodium chloride. It was in this way that Davy, in 1810, proved that com- mon salt is really nothing but sodium chloride. It combines directly with the other halogens. A molecule of sodium contains only one atom. Sodium is used in the laboratory to extract water from alcohol and ether and to prepare organic compounds. Large quantities are con- sumed in the manufacture of sodium peroxide (Na a O 2 ) and sodium cyanide (NaCN). Its power to reduce oxides gives it limited use in preparing certain metals, e.g. magnesium. Sodium Chloride, NaCl, is the most important compound of sodium. It is familiar under the name of salt or com- mon salt. The presence of salt in the ocean, in lakes and springs, and in the soil is mentioned in the oldest histori- cal records. It is one of the most abundant substances. The sources of salt are sea water, rock salt, and brines. Sodium, Potassium, and Lithium. 287 Preparation of Salt. Sea water contains nearly 4 per cent of salts, and three fourths of this amount is sodium chloride, (i) In warm countries, as on the shores of the Mediterranean Sea, shallow ponds of sea water near the shore are evaporated by exposure to the sun and wind, and the salt is collected. (2) In some regions sea water is first concentrated by allowing it to trickle over heaps of brush and then evaporated to crystallization in shallow pans. (3) In cold countries, as on the shores of the White Sea in Russia, sea water is allowed to freeze and the ice is removed. The ice contains no salt, so the opera- tion is repeated until the remaining liquid becomes strong enough to evaporate profitably over a fire. (4) Deposits of salt are found in many parts of the globe, the most important being in England, Austria- Hungary, and Germany. In these regions and some parts of the United States, the salt is mined and purified like other minerals. This variety is coarse and often impure, and is largely used in curing meat and preserving hides. (5) Most of the salt produced in the United States is obtained from natural or artificial brines, i.e. from strong solu- tions of salt. Artificial brines are made by forcing water into salt de- posits. Brines are obtained in New York, Michigan, Kansas, Ohio, West Virginia, California, Utah, and Louisiana. They are evaporated in vats by the sun's heat or by heating in kettles or pans. All these methods give a product containing as impurities salts of sodium, calcium, and magnesium, which are largely removed by further special treatment. The dampness of salt is due mainly to the magne- sium chloride it contains (see Deliquescence, Chapter IV). * Properties and Uses of Salt. Salt is soluble in water, 100 gm. of water dissolving about 36 gm. of salt at o C., and 40 gm. at 100 C. It crystallizes in cubes. This sub- stance is an essential ingredient of the food of man and animals. Besides its universal domestic use, enormous quantities are consumed in the preparation of many so- dium compounds, particularly sodium carbonate (see below), of hydrochloric acid and bleaching powder. In 1902 the United States produced nearly 3,000,000 tons of salt, and imported over 200,000 tons. This is about the average consumption. 288 Descriptive Chemistry. Sodium Carbonate, Na 2 CO 3 , is next to sodium chloride in importance. Small quantities of hydrated sodium car- bonates are found in Egypt, Russia, and in California and Nevada. Formerly it was obtained from the ashes of marine plants, but sodium chloride is now the source. The manufacture of sodium carbonate is one of the most extensive chemical industries. Two processes are used, the Leblanc and the Solvay. The Leblanc Process has three steps, (i) Sodium chloride is changed into sodium sulphate by sulphuric acid, the two equations for the changes being 2NaCl + H 2 SO 4 = HNaSO 4 + HC1 + NaCl Sodium Sulphuric Acid Sodium Hydrochloric Sodium Chloride Acid Sulphate Acid Chloride HNaSO 4 + NaCl Na,SO 4 + HC1 Sodium Sulphate This operation is called the "salt cake process 11 ; the impure prod- uct, called " salt cake," contains about 95 per cent of sodium sulphate. The hydrochloric acid is a by-product (see Hydrochloric Acid) . (2) The sodium sulphate is changed into sodium carbonate by heating the "salt cake" with coal and limestone, the main changes being repre- sented by the equations Na 2 S0 4 + 2 C Na 2 S + 2 CO, Sodium Sulphate Carbon Sodium Sulphide Carbon Dioxide Na 2 S + CaCO 3 Na 2 CO 3 + CaS Sodium Lime- Sodium Calcium Sulphide stone Carbonate Sulphide This operation is called the "black ash -process." The product is a dark brown or gray porous mass, and contains, besides 37 to 45 per cent of sodium carbonate, considerable calcium sulphide and other impuri- ties. The calcium sulphide is a source of sulphur (see Sulphur). (3) The sodium carbonate is rapidly separated from the insoluble portions of the " black ash " by agitation with a small amount of cool water. The solution of sodium carbonate thus obtained is evaporated to crys- Sodium, Potassium, and Lithium. 289 tallization, and the crude crystals are ignited. This product is known as soda ash, and from its solution in waiter are obtained soda crystals or sal soda (Na 2 CO 3 . 10 H,O). The Solvay Process, often called the ammonia -soda process, con- sists in saturating a cold concentrated solution of sodium chloride first with ammonia gas and then with carbon dioxide gas. The equation for the chemical change is . it NaCl + NH 3 + CO 2 = HNaCO., + NH 4 C1 Sodium Ammonia Carbon Acid Sodium Ammonium Chloride Dioxide Carbonate Chloride The acid sodium carbonate is nearly insoluble in the cold ammonium chloride solution, and therefore separates. It is changed, by heating, into sodium carbonate, thus 2 HNaCO, Na 2 CO 3 + CO 2 + . H 2 O Acid Sodium Sodium Carbon Water Carbonate Carbonate Dioxide The liberated carbon dioxide is used again, and from the ammonium chloride the ammonia is recovered and also used. Properties and Uses of Sodium Carbonate. Crystal- lized sodium carbonate (Na 2 CO 3 . 10 H 2 O) is often called alkali or soda. It loses water in the air, becoming dull at first and finally falling to a powder. When heated, it melts in its water of crystallization, and continued heating changes it into the white anhydrous salt (Na 2 CO 3 ). It is readily soluble in water, and the solution, which is strongly alkaline, is widely used as a cleansing agent, hence the name washing soda. Enormous quantities of sodium carbonate are used in the glass and soap industries, and in preparing sodium compounds. Sodium Bicarbonate, HNaCO 3 , is a by-product of the Solvay process, and it may also be prepared by treating crystallized sodium carbonate with carbon dioxide gas. It is a white powder, less soluble in water than the normal 290 Descriptive Chemistry. carbonate. When heated or when mixed with an acid or an acid salt, sodium bicarbonate gives up carbon dioxide. This property early led to its use in cooking, and gives the names cooking soda, baking soda, or simply soda. Sodium bicarbonate is one ingredient of baking powder and of the various mixtures (except yeast) used to raise bread, cake, and other food. Since cream of tartar is slightly acid,- it is usually used to liber- ate the gas. Sour milk, which contains lactic acid, is. sometimes used in place of cream of tartar. When pastry is raised with soda and cream of tartar, the escaping carbon dioxide puffs up the dough. Hence bak- ing soda is often called saleratus the salt which aerates (from the Latin words sal, salt, and aer, air or gas). Effervescing powders, such as Seidlitz (or Rochelle) and soda powders, contain sodium bicarbon- ate in one paper and tartaric acid or one of its acid salts in the other. When these are mixed in water, carbon dioxide is liberated. Sodium bicarbonate is used as a medicine to neutralize an acid stomach. For example, the " soda mints " sometimes taken for this purpose are mainly sodium bicarbonate. Sodium Hydroxide or Caustic Soda, NaOH, is a white corrosive solid. It absorbs water and carbon dioxide rapidly from the air. It dissolves readily in water, with rise of temperature, and the solution is strongly alkaline. It melts easily, and is often cast into sticks for use in the laboratory. Immense quantities are used in making hard soap, paper, and dyestuffs ; in bleaching, and in refining kerosene oil. Sodium hydroxide is usually manufactured by treating crude sodium carbonate with calcium hydroxide. Lime is added to a boiling, dilute solution of soda ash, and the main change is represented thus Ca(OH) 2 + Na 2 CO 3 = 2 NaOH + CaCO 3 Calcium Sodium Sodium Calcium Hydroxide Carbonate Hydroxide Carbonate The solution of sodium hydroxide is separated from the insoluble cal- cium carbonate, and concentrated by heating in iron kettles to the de- Sodium, Potassium, and Lithium. 291 sired strength or until the mass becomes stiff. Air is then blown in or sodium nitrate added to oxidize sulphides to sulphates. After standing several hours to allow other impurities to settle, the caustic soda is put into iron barrels called drums. It solidifies on cooling, and the drums are at once sealed to keep out the air. Sodium hydroxide is also manufactured on a large scale at Niagara Falls, New York, by the electrolysis of sodium chloride, according to the equation NaCl + Sodium Chloride H 2 O = NaOH + Sodium Hydroxide Cl Chlorine + H Hydrogen The apparatus is shown in Figure 56. The carbon anodes (A, A) pass into the outer compartments which contain brine, and the iron cathodes into the middle compartment which contains sodium hydroxide solution. When the cur- FlG. 56. Apparatus for the manufacture of sodium hydroxide by the electrolysis of sodium chloride. rent passes, chlorine is evolved at the anodes and flows out through pipes (not shown), and sodium is produced on the surface of the mercury (M ) which covers the floor of the whole apparatus. The sodium forms an amalgam 292 Descriptive Chemistry. with the mercury, and by rocking the apparatus on the device, B, B, the sodium amalgam flows into the compart- ment, D, where the sodium is liberated by the action of the electric current, which passes between the cathode and the amalgam. The sodium reacts with the water forming hydrogen, which passes off through pipes (not shown) and sodium hydroxide, which flows into a special tank. Both the chlorine and sodium hydroxide are nearly pure. The solution of caustic soda is finally treated, if necessary, as in the older process. Sodium Sulphate, Na 2 SO 4 , is one of the products obtained in the manufacture of sodium carbonate (see above). In another method, sulphur dioxide, steam, and air are passed into hot sodium chloride. And at Stassfurt, magnesium sulphate and sodium chloride are allowed to interact in the cold, thus MgSO 4 + 2NaCl = Na.SO 4 + MgCl 2 Magnesium Sodium Sodium Magnesium Sulphate Chloride Sulphate Chloride Sodium sulphate is a white anhydrous solid. It dissolves readily in water, and when a strong solution made at 30 C. is cooled, large transparent bitter crystals separate. They have the formula Na 2 SO 4 . ioH 2 O and are called Glau- ber's salt, from the discoverer. They lose water when exposed to air, and the salt continues to effloresce until it becomes an anhydrous powder. The crude salt is used in the glass and dyeing industries, and the purified salt as a medicine. Sodium Nitrate, NaNO 3 , is found abundantly in Chili, and is often called Chili saltpeter. It is a white solid, which becomes moist in the air. Large quantities are used as a fertilizer, either alone or mixed with compounds of Sodium, Potassium, and Lithium. 293 potassium and of phosphorus, and for making nitric acid and potassium nitrate. The natural deposits are in a dry region near the coast and cover over 200,000 acres. Chili controls the industry, and exports annually over a million tons. The crude salt, which looks like rock salt, is puri- fied by crystallization into a product containing 94-98 per cent of the nitrate. The final mother liquor is a source of iodine (see Iodine). Sodium Dioxide or Peroxide, Na.,O a , is a yellowish solid. It is used to bleach straw and delicate fabrics. With water it liberates oxygen, according to the equation Na 2 O 2 + H 2 O + 2NaOH Sodium Dioxide Oxygen Sodium Hydroxide Miscellaneous. Sodium cyanide (NaCN) is used to extract gold from poor ores. Sodium monoxide (Na 2 O) is a grayish solid. The sodium phosphates, sodium thiosulphate, acid sodium sulphite, sodium silicate, and sodium tetraborate or borax have been described. POTASSIUM. Occurrence. This metal is not found free, but its com- pounds are abundant. The minerals mica and feldspar are silicates containing potassium. By the decay of these and other minerals, potassium compounds find their way into the soil, thence into plants and animals. Potassium salts are found in wood ashes, in suint, the oily substance washed from sheep's wool, in beet-sugar residues, and in the deposits in wine casks. Sea water and mineral waters contain potassium salts, particularly potassium chloride and potassium sulphate. Many potassium salts are found at Stassfurt. About 2.5 per cent of the earth's crust is potassium. The Stassfurt deposits of the salts of potassium and other metals are near Magdeburg, Germany. About 16 different salts make up the beds, which are nearly 3000 feet thick. The deposits were doubt- less formed by the evaporation of sea water, though the different simpler 294 Descriptive Chemistry. salts interacted, forming complex ones. The most important salts Kainite .... KC1, MgSO 4 . 3 H 2 O. Carnallite . . . KC1, MgCl,, . 6 H 2 O. Polyhalite . . . K 2 SO 4 , Mg~SO 4 , 2 CaSO 4 . 2 H 2 O. Sylvite .... KC1. . Picromerite . . K 2 SO 4 , MgSO 4 . 6H 2 O. The name potassium comes from the word potash. The symbol, K, is from kalium, the Latin equivalent of kali, which is derived from an Arabic term for an alkaline substance. Preparation. Potassium is now obtained by the electrolysis of potassium hydroxide. Formerly it was manufactured, like sodium, by heating to a high temperature a mixture of potassium carbonate and carbon or of potassium hydroxide and iron carbide (see under Sodium). Properties. Like sodium, potassium is a soft, silver- white metal, light enough to float upon water the specific gravity being 0.86. Its brilliant luster soon disappears in air, owing to rapid oxidation. Potassium as ordinarily seen is, therefore, covered with a grayish coating, and, like sodium, must be kept under mineral oil. It melts at 62.5 C, and at a higher temperature burns with a violet-colored flame. This color is characteristic of burning potassium, arid is a test for the metal and its compounds. Like sodium, it decomposes water at ordinary temperatures, though more energetically. The heat evolved immediately ignites the hydrogen, and the melted potassium surrounded by a violet flame dashes to and fro upon the cold water. The main reaction corresponds to the equation K -f H 2 = KOH + H Potassium Water Potassium Hydroxide Hydrogen Potassium combines with the halogens and other ele- ments more vigorously than sodium, and forms analogous compounds. Sodium, Potassium, and Lithium. 295 Potassium Chloride, KC1, is found native in the Stass- furt deposits. It is also obtained in large quantities by decomposing carnallite and crystallizing the potassium chloride from the more soluble magnesium chloride. It is a white solid which crystallizes in cubes and otherwise resembles sodium chloride. It is used chiefly to prepare other potassium salts, especially the nitrate and chlorate. Potassium bromide and potassium iodide have been described (see Chapter XVI). Potassium Nitrate, KNO 3 , is also called niter and salt- peter. It is formed in the soil of many warm countries by the decomposition of nitrogenous organic matter (see Nitrification). It is now made by mixing hot, concentrated solutions of native so- dium nitrate and potassium chloride, which interact thus NaNO 3 + KC1 KN0 3 + NaCl Sodium Potassium Potassium Sodium Nitrate Chloride Nitrate Chloride The sodium chloride, being less soluble, separates, and is removed. By evaporation, small crystals of potassium nitrate, called " niter meal, 11 are obtained, and further purified by recrystallization. Potassium nitrate is a white solid. It dissolves easily in cold water with a fall of temperature, and very freely in hot water, but it is not hygroscopic. It is crystalline, but con- tains no water of crystallization. The taste is salty and cooling. It melts at 339 C., and further heating changes it into potassium nitrite (KNO 2 ) and oxygen. At a high temperature, potassium nitrate gives up oxygen readily, especially to charcoal, sulphur, and organic matter. This oxidizing power leads to its extensive use in making gun- powder, fireworks, matches, explosives, and in many chemi- cal operations. 296 Descriptive Chemistry. Gunpowder is a mixture of potassium nitrate, charcoal, and sulphur. The ingredients are first purified, pulverized, and thoroughly mixed. This mixture is pressed, while damp, into a thin sheet ; and the " press cake" thus formed is broken into small grains, which are sorted by sieves. The grains are then smoothed or ''glazed" by rolling them in a barrel, again sifted, arid finally dried at a low temperature. The pro- portions differ with the use of the powder. The United States army standard black powder contains 75 per cent of potassium nitrate, 15 of charcoal, and 10 of sulphur. When gunpowder burns in a closed space, a large volume of gas is suddenly formed. So enormously is this gas expanded by the heat that it would fill several hundred times the space taken by the powder itself. The pressure exerted by this expanding gas is many tons. It is this pressure which forces the ball from a cannon and tears a rock to pieces. The chemical changes attending the explo- sion of gunpowder in a closed space are complex, as may be seen by the following (approximate) equation : 8 KNO 3 + 90 + 38 = 2 K,C0 3 + K 2 SO 4 -f K 2 S., + 7 CO 2 + 4 N 2 Probably secondary reactions produce other gases besides carbon diox- ide and nitrogen. Potassium Chlorate, KC1O 3 , is a white, crystallized, lus- trous solid. It tastes like potassium nitrate. It melts at 334 C., and at a high temperature decomposes into oxygen and potassium chloride as final products, thus KC1O 3 KC1 + O 3 Potassium Chlorate Potassium Chloride Oxygen It is used to prepare oxygen, and in the manufacture of matches and fireworks. In the form of " chlorate of potash tablets " it is used as a remedy for sore throat. Potassium chlorate is manufactured by passing chlorine into calcium hydroxide (milk of lime) and adding potassium chloride to the mixture. The simplest equations for the complex changes may be written thus : (i) 6 Ca(OH) 2 + 6 C1 2 = Ca(ClO 3 ) 2 + 5 CaCl 2 + 6 H,O Calcium Calcium Calcium Hydroxide Chlorate Chloride Sodium, Potassium, and Lithium. 297 (2) 3 Ca(ClO) 2 Ca(C10 3 ) 2 + 2 CaCl 2 Calcium Hypochlorite Calcium Chlorate (3) Ca(ClO 3 ) 2 + 2 KC1 2 KC1O 3 + CaCl 2 Potassium Chlorate The salt is also made by the electrolysis of a hot solution of potassium chloride, though it has been found more satisfactory to first prepare sodium chlorate and convert this salt into potassium chlorate by po- tassium chloride. Potassium Carbonate, K 2 CO 3 , is a white powder. It deliquesces in the air, is very soluble in water, and the solution has a strong alkaline reaction. It was formerly obtained by treating wood ashes with water, and evaporating the solution to dryness. The crude salt thus obtained has long been called potash, and a purer product is known as pearlash. (The term potash is sometimes applied to potassium oxide, K 2 O.) It is used extensively in the manu- facture of hard glass, soft soap, caustic potash, and other potassium compounds. Potassium carbonate is obtained from suint by igniting the greasy mass and extracting the potassium carbonate with water. Beet-sugar residues also furnish potassium carbonate. After the sugar has been obtained from the beet sirup, the molasses is changed by fermentation into alcohol, which is distilled off; the liquid residue is evaporated to dryness and ignited, and the potassium carbonate extracted with water. Pure potassium carbonate is prepared by igniting cream of tartar made from the deposits in wine casks. All these sources emphasize the inti- mate relation of potassium compounds to vegetable and animal life. The bulk of the potassium carbonate is now made from potassium sul- phate or from the chloride by the Leblanc process, owing to the abun- dance of crude potassium salts at Stassfurt. Potassium Hydroxide or Caustic Potash, KOH, is a white brittle solid, resembling caustic soda. It absorbs water and carbon dioxide very readily ; and if exposed to the air, soon becomes a thick solution of potassium 298 Descriptive Chemistry. carbonate. Like sodium hydroxide, it dissolves in water with evolution of heat, forming a strongly alkaline caustic solution. It is one of the strongest bases, even glass and porcelain being corroded by it. Besides its use in the labo- ratory, large quantities are consumed in making soft soap. Potassium hydroxide is made and purified in the same way as sodium hydroxide, viz. by adding lime or milk of lime to a boiling dilute solution of potassium carbonate, the equation for the change being : Ca(OH) 2 -f K,CO 3 2 KOH + CaCO 3 Milk of Potassium Potassium Calcium Lime Carbonate Hydroxide Carbonate It is also made by the electrolysis of a solution of potassium chloride. Miscellaneous. Potassium Cyanide (KCN) is a white solid, very poisonous, very soluble in water, and having an odor like bitter almonds (see Cyanogen, Chapter XIV) . Potassium Sulphate (K 2 SO 4 ) is manu- factured from kainite, and is largely used as a fertilizer and in making potassium carbonate. Relation of Potassium to Life. Potassium, like nitro- gen and phosphorus, is essential to the life of plants and animals. The ash of many common grains, vegetables, and fruits contains potassium as the carbonate. Potassium salts are supposed to assist in the formation of starch, just as phosphorus is indispensable to the transformation of nitro- gen compounds. Potassium salts taken from the soil by plants must be returned if the soil is to be productive. Sometimes crude kainite is used extensively as a fertilizer ; but wood ashes, or the sulphate and chloride, are often used to supply potassium salts. Lithium, Li, is a silver-white metal and has the specific gravity of only 0.59. It is the lightest of the metallic elements. Its compounds are widely distributed in small quantities in minerals, mineral waters, and plants. Lithia water and citrate of lithium are often prescribed as a remedy for diseases of the kidneys. Lithium compounds color the Bunsen flame bright red a delicate test for the metal. Sodium, Potassium, and Lithium. 299 Rubidium and Caesium, Rb and Cs, have properties and form com- pounds analogous to those of potassium. EXERCISES. 1. Name the alkali metals. What is the symbol of each ? When and by whom was each discovered ? 2. What are the important compounds of sodium ? What per cent of the earth's crust is sodium ? 3. Describe the manufacture of sodium by electrolysis. Describe the older methods of manufacture. 4. Summarize (#) the physical properties, and (<) the chemical properties of sodium. How is it usually kept ? For what is it used ? 5. Discuss the interaction of sodium and water (see Chapter V). 6. Give the chemical name and formula of common salt Where is it found ? 7. Describe the different methods of preparing salt. State (#) the properties, and (b) the uses of salt. 8. Discuss the manufacture of sodium carbonate by (a) the Le- blanc process, (b) By the Solvay process. 9. What is (a) soda, (b) soda ash, (V) salt cake, (-. Calcium, Strontium, and Barium. 329 making fireworks, especially "red fire." The latter is a mixture of potassium chlorate, shellac, and strontium nitrate. The production of the crimson colored flame is the test for stron- tium. Compounds of Barium. The most abundant native compounds are witherite (barium carbonate, BaCO 3 ) and barite (barium sulphate, BaSO 4 ) . The oxides, BaO and BaO 2 , have already been mentioned as a source of oxygen. Barium hydroxide (Ba(OH) 2 ) solution is often called baryta water, and it forms the insoluble barium carbonate (BaCO 3 ) when exposed to carbon dioxide. Barium chloride (BaCl 2 ) is used in the laboratory to test for sulphuric acid and soluble sulphates, because it readily interacts with them and forms the insoluble barium sulphate (BaSO 4 ). This precipitated salt is a fine, white powder, and being cheap and heavy it is a common adulterant of the ordinary white paint, Ground native barium sulphate has a similar use. Barium sul- phate is also used to increase the weight of paper and to give it a gloss. Barium salts color a flame green, and barium nitrate (Ba(NO 3 ) 2 ) is extensively used in making fireworks, especially "green fire." Com- mercial barium sulphide (BaS), as well as the sulphides of calcium and strontium, shine feebly in the dark, after having been exposed to a bright light. On account of this property they are used in making luminous paint. Soluble barium salts are poisonous. The production of the green flame is the test for barium. EXERCISES. 1. Name the alkaline earth metals. What is the symbol of each ? 2. Name several compounds of calcium. What proportion of the earth's crust is calcium ? 3. Describe the preparation and state the properties of calcium. 4. What is the formula of calcium carbonate ? State the properties, occurrence, and uses of (a) limestone, and () marble. 5. State the essential characteristics of (a) calcite, () Iceland spar, (c) stalactites, (d) Mexican onyx, (e) travertine, (/) coquina, (g) chalk, (Ji) coral. 6. Review the properties of calcium carbonate, especially its solu- bility (see Chapter XIV). 7. State the uses of (a) limestone, () marble, (c) chalk. 8. Describe the formation of (a) limestone caves, () chalk, (<:) coral. jjo Descriptive Chemistry. 9. What is the formula and chemical name of lime ? State the properties and uses of lime. How is it made ? State the equation for the chemical change. 10. What is (a) quicklime, (fr) slaked lime, (c) hydraulic lime, (d) Portland cement, (e) " air-slaked " lime ? 1 1 . What is the formula of calcium hydroxide ? How is it formed ? What are its properties ? How does it interact with carbon dioxide ? State the equation for the reaction. 12. What is () limewater, (b) milk of lime, (c~) whitewash ? 13. What is mortar ? How is it prepared ? For what is it used ? How does it change chemically with age ? What is plaster ? 14. Discuss the occurrence of calcium sulphate. State the chief properties of (a} gypsam, (b} selenite, (c) alabaster, (d) satin spar. For what are gypsum and alabaster used ? 15. What is plaster of Paris ? Why so called ? How is it pre- pared ? What is its chief property ? What are its uses ? What is the chemical explanation of " setting '' ? What is stugco ? 1 6. What is hard water ? How does it act with soap ? What is (#) temporary hardness, and (b) permanent hardness ? How may each be removed ? What is soft water ? Why is rain water often called soft water ? 17. Summarize the properties of calcium chloride. What is its formula ? How is it prepared ? 1 8. Review the essential properties of (a) calcium fluoride, (b) cal- cium carbide, (c) tricalcium phosphate, (d} bleaching powder. 19. What is the test for (a) calcium, (b) strontium, (c) barium ? 20. State the use of (a) strontium hydroxide, and {b) strontium nitrate. 21. For what are (a) barium hydroxide, (b} barium nitrate, (V) barium sulphide, and (//) barium chloride used ? Describe barium sulphate. PROBLEMS. 1. What is the per cent of calcium in (a) marble (CaCO <3 ), () gypsum (CaSO 4 . 2 H.,O), (V) fluor spar (CaF 2 ), (d} superphos- phate of lime (CaH 4 (PO 4 ) 2 ) ? 2. How many tons of limestone must be heated to produce 100 tons of quicklime ? (Assume CaCO 3 = CaO + CO 2 .) 3. Calculate the simplest formula of a compound having the per- centage composition Ca = 40, C = 12, O = 48. CHAPTER XXIV. MAGNESIUM, ZINC, CADMIUM, AND MERCURY. THESE elements form a natural group, though the mem- bers are not so closely related as the alkali and alkaline earth groups. Zinc and cadmium are much alike, and both also resemble magnesium. Mercury differs somewhat from zinc and cadmium, but resembles copper. MAGNESIUM. Occurrence of Magnesium. Magnesium is never found free. In combination it is widely distributed and very abundant, constituting about 2.5 per cent of the earth's crust. Dolomite is magnesium calcium carbonate (CaMg(CO 3 ) 2 ) ; it forms whole mountain ranges and vast deposits; beds hundreds of feet thick cover thousands of square miles in the upper Mississippi valley. Dolomite closely resembles marble and limestone. Magnesium carbonate is also abundant. Many of the Stassfurt salts contain magnesium, for example, kainite (KC1, MgSO 4 . 3 H 2 O), carnallite (KC1, MgCl 2 .6 H 2 O), and kieserite (MgSO 4 . H 2 O). It is also a component of serpentine, talc, soapstone, asbestos, meerschaum, and other silicates. The sulphate and chloride are found in sea water and in mineral springs. Through the decay of rocks, magnesium compounds find their way into the soil, from which they are taken up by plants. Magnesium phosphates are found in the bones of animals and the seeds of grains, and also in guano, Descriptive Chemistry. D II (7= Preparation of Magnesium. Magnesium was formerly prepared by reducing the chloride with sodium. It is now economically manufac- tured by electrolysis. A sketch of the essential parts of the apparatus is shown in Figure 67. Carnallite is put into the cylindrical iron vessel, C, which is the cathode. This is closed by the air-tight cover through which pass the pipes, Z>, D ', for conveying inert gases into and out of the apparatus. The carbon anode, A, dips into the carnallite and is inclosed by the porcelain cylinder, B, which is provided with a pipe, E, for the escape of the chlorine liberated at the anode. The carnallite is kept fused by external heat. When the current passes, the chlorine liberated at the anode escapes through E, and the magnesium liberated at the cathode floats on the fused carnallite and is pre- vented from oxidizing by the inert gas supplied through D. The porcelain cylinder, B, prevents the chlorine from escaping into the larger vessel. The FIG. 67. Apparatus 'for the manufacture of magnesium by the molten magnesium is carefully removed electrolysis of carnallite. at intervals. Properties of Magnesium. Magnesium is a lustrous, silvery white metal. It is a light metal, the specific grav- ity being only 1.75. It is tenacious and ductile, and when hot may be drawn into wire or pressed into ribbon, the latter being a common commercial form. It melts at a red heat and may be cast into different shapes. At a high temperature it volatilizes. It is easily kindled by a match or candle, and burns with a dazzling white light, producing dense white clouds of magnesium oxide (MgO). It does not tarnish in dry air, but in moist air it is soon covered with a film of oxide. It liberates hydrogen from acids. Heated in nitrogen, it forms magnesium nitride (Mg 3 N 2 , see Composition of Ammonia). Uses of Magnesium. Magnesium in the form of pow- der is used chiefly in taking flash-light photographs. Magnesium, Zinc, Cadmium, and Mercury. 333 Small quantities are used in making fire-works ; and both the powder and wire are used in the chemical laboratory. Magnesium Oxide, MgO, is a white, bulky powder. It is formed when magnesium burns in the air, but it is man- ufactured by gently heating magnesium carbonate, just as lime is made from limestone. It is often called magnesia, or calcined magnesia. The native oxide is the mineral periclase. Magnesia dissolves with difficulty in water, forming magnesium hydroxide (Mg(OH) 2 ). A mixture of magnesia and water, with or without magnesium chlo- ride, hardens on exposure to the air, and is often used as a cement or artificial stone. Native magnesium hydroxide is the mineral brucite. Like lime, magnesia withstands a high temperature, and is, therefore, used as the chief ingredient of a protective mixture for steam pipes and ves- sels which are subjected to great heat. Magnesia is used as a medicine for dyspepsia and an antidote for poisoning by mineral acids. Magnesium Sulphate, MgSO 4 , is a white solid. There are several crystallized varieties. The native salt kie- serite (MgSO 4 . H 2 O) when added to water changes into Epsom salts (MgSO 4 . ;H 2 O). This variety was first found in the mineral spring at Epsom, England. It is very soluble in water, and its solution has a bitter taste. It is extensively used as a medicine, in manufacturing sul- phates of sodium and potassium, as a fertilizer in place of gypsum, and as a coating for cotton cloth. Magnesium Chloride, MgCl 2 , is a white solid. It is a by- product in the preparation of potassium chloride. The crystallized salt (MgCl 2 . 6 H.,O) is very deliquescent. Magnesia mixture is a mix- ture of magnesium chloride, ammonium chloride, and ammonium hy- droxide ; it is used in chemical analysis. 334 Descriptive Chemistry. Magnesium Carbonate, MgCO 3 , occurs native as magnesite, and combined with calcium carbonate as dolomite. The commercial salt known as magnesia alba, or simply magnesia, is a complex compound (Mg(OH) 2 , 4 MgCO 3 4 H 2 O) . Several of these complex basic carbon- ates are known. Many face powders consist chiefly of magnesia alba. It was during an investigation of magnesia alba that Black discov- ered carbon dioxide and showed the close relation between analogous compounds of magnesium and calcium. Miscellaneous. Besides the oxide and sulphate, other compounds are used as medicines. Fluid magnesia, prepared by dissolving mag- nesium carbonate in water containing carbon dioxide, is a mild laxative. Magnesium citrate has a similar action ; it is an effervescing mixture prepared from sodium bicarbonate, tartaric and citric acids, sugar, and magnesium sulphate. ZINC. Occurrence of Zinc. Free zinc is never found. The ores of zinc are not numerous, but are widely distributed. The chief ores are zinc sulphide (sphalerite, zinc blende, ZnS), zinc carbonate (smithsonite, ZnCO 3 ), zinc silicate (calamine, H 2 Zn 2 SiO 5 ), and red zinc oxide (zincite, ZnO). Franklinite and willemite are ores of zinc containing manganese and iron. Gahnite has the composition ZnAl 2 O 4 . Zinc ores are found in Germany, Italy, France, Greece, Spain, Austria- Hungary, Belgium, England, and the United States. Missouri and Kansas contain large deposits of the sulphide, while the other ores occur chiefly in New Jersey. About 143,000 tons of zinc were pro- duced in the United States in 1902, and over 60 per cent came from Missouri-Kansas. This was the largest amount ever produced in a single year. Metallurgy of Zinc. Zinc is easily smelted. The ores are first roasted to change them into the oxide, thus ZnCO 3 = ZnO + CO 2 Zinc Carbonate Zinc Oxide Carbon Dioxide Magnesium, Zinc, Cadmium, and Mercury. 335 ZnS +30= ZnO + SO 2 Zinc Sulphide Oxygen Zinc Oxide Sulphur Dioxide The oxide is then reduced by heating it with charcoal. This operation is conducted in earthenware tubes or fire- clay crucibles connected with iron receivers into which the zinc vapor passes ; at first it condenses as a powder known as zinc dust, somewhat as sulphur forms flowers of sul- phur ; but it finally condenses as a liquid, which is drawn off at intervals and cast into bars or plates. The impure zinc thus obtained is called spelter ; it is freed from carbon, lead, iron, cadmium, and arsenic by repeated distillation, often under reduced pressure. Properties of Zinc. Zinc is a bluish white, lustrous metal. Its physical properties vary with the temperature. At ordinary temperatures it is brittle, but at 100 150 C. it is soft and may be rolled into sheets and drawn into wire, while its specific gravity rises from 6.9 to 7.2. Zinc which has been rolled or drawn does not become brittle upon cooling. At 200 C. it again becomes brittle and can be easily pulverized. It melts at about 433 C. and boils at about 940 C. Heated in the air above its melting point, zinc burns with a bluish green flame, forming white zinc oxide (ZnO). Zinc does not tarnish in dry air, but ordinarily it becomes coated with a dark film. Commercial v zinc interacts with acids and usually liberates hydrogen. With hot solutions of sodium and potassium hydroxides, it forms zincates and liberates hydrogen, thus 2KOH + Zn = H 2 + K 2 ZnO 2 . Potassium Hydroxide Zinc Hydrogen Potassium Zincate Pure zinc interacts with acids if in contact with a platinum wire, or if copper sulphate solution is added. Like copper, 33 6 Descriptive Chemistry. zinc withdraws other metals (e.g. lead and mercury) from their solutions. The vapor density of zinc requires the molecular weight 67.6. Since the atomic weight is 65.4, a molecule of the vapor contains only one atom. Uses of Zinc. Zinc in stick or plates is extensively used as the positive plate in electric batteries. Sheet zinc is used as a lining for tanks, and as the protective cover- ing which is placed behind and beneath stoves. Iron dipped into melted zinc becomes coated with zinc and is called galvanized iron ; it does not rust easily and is widely used for roofs, pipes, cornices, and water tanks. Telegraph wire is also galvanized. Zinc dust is used in the cyanide process of extracting gold and in many chemical experi- ments in the laboratory. Brass, German silver, and other alloys contain zinc (see Alloys of Copper). Antifriction metals, which are used for bearings, are alloys of zinc. Babbitt's metal, for example, contains 69 per cent of zinc, 19 of tin, 4 of copper, 3 of antimony, and 5 of lead. Compounds of Zinc. Native zinc oxide is red, owing to the presence of manganese, but the pure oxide is white when cold and yellow when hot. It is formed when zinc burns, and is manufactured in this way or by heating zinc carbonate. It is often called "zinc white" or " Chinese white," and is used to make a white paint which is not dis- colored by the atmosphere. Native zinc sulphide is yel- low, brown, or black on account of impurities, but the pure sulphide is white. The latter is formed as a jelly like pre- cipitate when hydrogen sulphide is passed into an alkaline solution of a zinc salt ; it is decomposed by a mineral acid. Zinc sulphide is also used as a white pigment. Zinc sulphate is. formed by the interaction of zinc and dilute sulphuric acid. Large quantities are made by roasting Magnesium, Zinc, Cadmium, and Mercury. 337 the sulphide in a limited supply of oxygen and extracting the sulphate with water. It is a white, crystallized solid (ZnSO 4 . 7 H 2 O), which effloresces in the air, and when heated to 100 C. loses most of its water of crystallization. The crystallized salt is called white vitriol. It is used in dyeing and calico printing, as a disinfectant, and as a medi- cine. It is poisonous, but can be safely used externally to relieve inflammation. Zinc chloride (ZnCl 2 ) is a white, deliquescent solid, prepared by dissolving zinc in hydro- chloric acid and evaporating the solution until a sample solidifies on cooling. It is used in surgery, and also as a constituent of a mixture for filling teeth ; large quantities are used to preserve wood, especially railroad ties, from decay, nearly 1500 tons being annually consumed for this purpose. Zinc hydroxide (Zn(OH) 2 ) is formed by the interaction of sodium or potassium hydroxide and the solu- tion of a zinc salt. An excess of the alkaline hydroxide changes the zinc hydroxide into a zincate. Tests for Zinc. The formation of the sulphide or hydroxide, as above described, serves as the test for zinc. A green incrustation is produced when zinc compounds are heated on charcoal and then mois- tened with a cobaltous nitrate solution. Cadmium, Cd, is an uncommon metal, frequently found in zinc ores. It occurs native as a sulphide (greenockite, CdS). It is white, lustrous, and rather soft. Its specific gravity is 8.6, and its melting point is about 320 C. Cadmium is a constituent of certain fusible alloys (see Bismuth). Wood's metal contains 12 per cent of cadmium. The most important compound is cadmium sulphide (CdS). This is a bright yellow solid, formed by adding hydrogen sulphide to the solution of a cadmium compound. It is used as an artist's color. Its formation also serves as the test for cadmium. MERCURY. Occurrence of Mercury. Native mercury is occasion- ally found in minute globules, but the most abundant ore jj 8 Descriptive Chemistry. is mercuric sulphide (cinnabar, HgS). The ore is mined in Spain, Austria, Russia, Italy, and Mexico ; in the United States large quantities are obtained in California, and deposits were recently opened in Texas. The annual production of the United States for several years has been about 1000 tons. Mercury has been known for ages as quicksilver. The Latin name, hydrargyrum, which gives us the symbol Hg, means literally " water silver," emphasizing the fact, so well known, that mercury looks like silver and flows like water. Preparation of Mercury. Mercury is readily prepared by roasting cinnabar in a current of air. Sulphur dioxide and mercury are formed, thus HgS + 2 Hg + S0 2 Cinnabar Oxygen Mercury Sulphur Dioxide The sulphur dioxide is usually allowed to escape, but the mercury vapor is condensed by passing it into large cham- bers, or through pear-shaped retorts or pipes, called aludels (see Iodine). Crude mercury is freed from dirt and me- chanical impurities by pressing it through linen or chamois leather, but it must be distilled to separate it from dissolved metals, such as lead or zinc. It can also be purified by treatment with dilute nitric acid. Mercury is sent into commerce in strong iron flasks, holding about 75 pounds. Properties of Mercury. Mercury is a bright, silvery metal, and is the only one which is liquid at ordinary tem- peratures. It solidifies at about 39.5 C. It is a heavy metal, the specific gravity being 13.59. It is slightly vola- tile even at ordinary temperatures, and the vapor is poison- ous. Mercury does not tarnish in the air, unless sulphur compounds are present. At a high temperature, it com- bines slowly with oxygen to form the red oxide (HgO). Magnesium, Zinc, Cadmium, and Mercury. 339 Hydrochloric acid and cold sulphuric acid do not affect it ; hot concentrated sulphuric acid oxidizes it, and nitric acid changes it into nitrates. The vapor density of mercury requires the molecular weight 198.72. Since the atomic weight is 200, a molecule of the vapor contains only one atom. Amalgams are alloys of mercury with other metals. They are easily prepared by mixing the constituents. Sometimes the union is violent as in the preparation of sodium amalgam. Amalgamated zinc is usually used in electric batteries to prevent unnecessary loss of the zinc. Tin amalgam is sometimes used to coat mirrors. Amal- gams of certain metals are used as a filling for teeth. Care should be taken, while handling mercury, not to let it come in contact with rings or jewelry, since gold amalgam is readily formed. Uses of Mercury. Mercury is used in making ther- mometers, barometers, and some kinds of air pumps. Its extensive use in extracting gold and silver has been men- tioned (see Amalgamation). Large quantities are used in preparing certain medicines and explosives (e.g. fulminating mercury, which is used in cartridges). Compounds of Mercury. Mercury, like copper, forms two classes of compounds the mercurous and the mercuric. Mercuric oxide (HgO) is a red powder, produced by heating mercury in air or by heating a mixture of mercury and mercuric nitrate. As we have already seen, mercuric oxide is decomposed by heat into mercury and oxygen. A yellow variety is produced by the interaction of sodium hydroxide and a mercuric salt, thus 2NaOH + Hg(NO 3 ) 2 = HgO + 2 NaNO 3 + H 2 O Sodium Mercuric Mercuric Sodium Hydroxide Nitrate Oxide Nitrate Mercurous chloride (Hg 2 Cl 2 or HgCl) is a white, tasteless powder, insoluble in water. It is formed when a chloride and mercurous nitrate 34-O Descriptive Chemistry. interact, but it is manufactured by heating a mixture of mercuric chloride and mercury. Under the name of calomel it is extensively used as a medicine. Mercuric chloride (HgCl 2 ) is a white, crystalline solid, solu- ble in water and in alcohol. It is prepared by heating a mixture of mercuric sulphate and common salt. It is a violent poison. The best antidote is the white of a raw egg. The albumen forms an insoluble mass with the poison, which may then be removed mechanically from the stomach. The common name of mercuric chloride is corrosive sublimate. It has strong antiseptic properties, and is extensively used in surgery to protect wounds from the harmful action of germs ; taxi- dermists sometimes use it to preserve -skins, and it has many serviceable applications as a medicine and disinfectant. It is usually used as a dilute solution (i part to 1000 parts of water). Native mercuric sul- phide or cinnabar (HgS) is a red, crystalline solid. When hydrogen sulphide is passed into a solution of a mercuric salt, mercuric sulphide is formed as a black powder; this variety, when heated, changes into red crystals. Vermilion is artificial mercuric sulphide. It is manufactured either (i) by grinding together mercury and sulphur, and treating this mass with caustic potash solution, or (2) by heating mercury and sulphur in iron pans and subliming the black mass. In both processes the product must be carefully ground, washed, and dried. Chinese vermilion is the best quality. Vermilion has a brilliant red color, and, although expen- sive, is widely used to make red paint. Mercurous Nitrate (HgNO 3 or Hg 2 (NO 3 ) 2 ) and mercuric nitrate (Hg(NO 3 ) 2 ) are prepared by treating mercury respectively with cold dilute nitric acid, and with hot concentrated nitric acid. They are white, crystalline solids. EXERCISES. 1. Name the chief native compounds of magnesium. What pro- portion of the earth's crust is magnesium ? 2. Describe the manufacture of magnesium by the electrolysis of carnallite. 3. Summarize the properties of magnesium. State its uses. 4. What is the formula and chemical name of magnesium ? How is magnesia formed ? State its properties and uses. 5. Describe the different varieties of magnesium sulphate. State the uses of Epsom salts. Magnesium, Zinc, Cadmium, and Mercury. 341 6. What is the formula of magnesium carbonate ? What is (#) magnesite, () dolomite, (c) magnesia alba? For what is the last sub- stance used ? 7. Name the chief ores of zinc. Discuss their occurrence. 8. Describe the metallurgy of zinc. What is (a) zinc dust, and () spelter ? How is zinc purified ? 9. Summarize (a) the physical properties of zinc, and () the chem- ical properties. 10. State the uses of zinc. 11. Review the alloys of copper which also contain zinc. What alloys are largely zinc ? 12. Describe native and pure zinc oxide. For what is the latter used ? 13. Describe zinc sulphate. How is it formed and for what is it used? 14. Describe zinc chloride. For what is it used? 15. What are the tests for zinc ? 1 6. State the properties and uses of (a) cadmium, and () cadmium sulphide. 17. What is the chief ore of mercury ? Where is it found ? 1 8. What is the symbol of mercury? What is the literal meaning of the word from which it is formed ? 19. Describe the preparation and purification of mercury. How is it transported ? 20. Summarize the properties of mercury. 21. What are amalgams ? Name three, and state the use of each. 22. For what is mercury used ? 23. Describe mercuric oxide. What historical interest has it ? 24. Describe mercurous chloride. What is its commercial name? State its use. 25. Describe mercuric chloride. What is its commercial name? How does it differ from mercurous chloride ? State its use. 26. What is the formula and chemical name of cinnabar ? Describe cinnabar. What is vermilion ? How is it manufactured? State its use. 27. What is (a) magnesia, () Epsom salts, (c) galvanized iron, (d) Chinese white, 0) white vitriol, (/) calomel, (g) corrosive subli- mate ? 342 Descriptive Chemistry. PROBLEMS. 1. How much magnesium will be formed by heating 100 gm. of potassium with magnesium chloride ? (Assume K 2 + MgCl 2 = Mg + 2 KC1.) 2. What is the per cent of magnesium in (#) magnesite (MgCO 3 ), () dolomite (MgCa(CO 3 ) 2 ), (c} Epsom salts (MgSO 4 7 H,O ) ? 3. What is the per cent of zinc in (a) zinc sulphate (ZnSO 4 ), (b} zinc sulphide (ZnS), (c) zinc chloride (ZnCl 2 ), (d) zinc oxide (ZnO) ? 4. How much zinc sulphate can be prepared from 65 gm. of zinc ? From 130 gm.? From 720 gm.? 5. How much mercury is formed by decomposing 400 gm. of cin- nabar ? (Assume HgS + O 2 = Hg + SO 2 .) 6. What is the per cent of mercury in (a) mercuric oxide (HgO), (b) calomel (Hg 2 Cl 2 ), (c) corrosive sublimate (HgCl 2 ) ? CHAPTER XXV. ALUMINIUM. Occurrence. Aluminium does not occur free in nature, but its compounds are numerous, abundant, and widely distributed. About 8 per cent of the earth's crust is aluminium; it is, therefore, the most abundant metal. Many common rocks and minerals are silicates of alumin- ium and other metals, e.g. feldspar and mica, which make up a large part of granite and gneiss. Clay and slate are mainly silicate of aluminium, formed by the decomposition of complex aluminium minerals. Corundum and emery are aluminium oxide (A1 2 O 3 ) more or less impure. Baux- ite is an hydroxide of aluminium (H 4 A1 2 O 5 ). Cryolite is a fluoride of aluminium and sodium (Na 3 AlF 6 ). Aluminium was first obtained as a fine powder by Wohler in 1827. Deville, in 1854, prepared it in compact form and laid the foundation of the industry which is being developed by Hall. Davy proposed the name alumium, i.e. alum + him, to emphasize the relation of the metal to the well-known substance, alum. The word alumium was changed first to aluminum and then to aluminium. Some authorities derive the word alumium from the Latin word alumen, or from alumina, the common name of aluminium oxide. Metallurgy. Aluminium is obtained from its oxide (A1 2 O 3 ) by electrolysis. In the Hall process, which is typical, an open, iron box lined with carbon is made the cathode (Fig. 68). The anode consists of carbon bars hung from a copper rod, which can be lowered as the car- 343 344 Descriptive Chemistry. bon is consumed. The process is essentially as follows : the bottom of the box is covered with cryolite, the anodes are lowered, and the box is then filled with cryolite. The current is turned on, and in its resisted passage through the cryolite enough heat is generated to melt the cryolite. Pure, dry aluminium oxide is now added. This is decom- R . posed into aluminium and oxygen. The oxy- gen unites with the carbon of the anodes, forming carbon mo- noxide, which burns or escapes. The molten FlG. 68. Apparatus for the manufacture of aluminium falls to the aluminium by the electrolysis of aluminium oxide. C C C is the iron box which serves as the cathode. A, A, etc. are carbon anodes attached to the copper rod, R. bottom. The process is continuous, fresh aluminium oxide being added and the molten aluminium being drawn off at inter- vals. The cryolite is unchanged, and merely acts as a solvent for the aluminium oxide. The United States produced about 7,000,000 pounds of aluminium in 1902, and the output is annually increasing. This was all produced at Niagara Falls. In the Heroult process, which is used in Europe and involves essentially the same principle as Hall's process, the aluminium is produced as an alloy (usually of copper) . Aluminium was prepared until about 1885 by a complicated process, (i) Bauxite was changed into aluminium oxide free from iron by fusion with sodium carbonate and treatment with carbon dioxide. (2) The aluminium oxide was then changed into aluminium sodium chloride by fusion with sodium chloride and charcoal and subsequent treatment with chlorine. (3) This chloride was reduced by sodium, thus A1C1 3 + 3Na = Al + 3 NaCl Aluminium Sodium Aluminium Sodium Chloride Chloride Aluminium. 345 The sodium for this operation was prepared by the Castner process (see Sodium), and the two industries were developed simultaneously. The extensive application of the electrolytic method has reduced the price of aluminium from about $12 a pound during 1862-1887 to about 30 cents in 1902. Properties. Aluminium is a bluish white metal. It is very light compared with other common metals, since its specific gravity is only about 2.6 ; this value is one third that of iron. It is ductile and malleable, and is often sold in the form of wire and sheets ; it must be annealed frequently during the hammering or drawing. It is a good conductor of heat and electricity. Its tensile strength is about as great as that of cast iron. It melts at about 660 C., and may be cast and welded, but not readily soldered so as to produce a permanent joint. The cap of the Washington Monument is a casting of aluminium which weighs about eight and a half pounds. Pure alu- minium is only very slightly oxidized by air. Hydrochlo- ric acid changes it into aluminium chloride, thus 2A1- + 6HC1 = 2A1C1 3 + 3H 2 Aluminium Hydrochloric Aluminium Hydrogen Acid . Chloride Under ordinary conditions nitric and' sulphuric acids do not affect it. Sodium and potassium hydroxides change it into aluminates, thus 6NaOH + 2A1 = 2 Na 3 AlO 3 + 3 H 2 Sodium Hydroxide Aluminium Sodium Alumkiate Hydrogen The properties of aluminium are modified by the presence of impuri- ties. The usual impurities are iron, other metals, and silicon. Some of these, especially the iron and silicon, come from the raw products used in its manufacture. They tend to make the metal harder and more active chemically, but less malleable, ductile, and tenacious. If it were not for the presence of these impurities in clay, this substance would be a cheap and inexhaustible source of aluminium. 346 Descriptive Chemistry. Uses. The varied properties of aluminium adapt it to numerous uses. It is made into the metallic parts of mili- tary outfits, caps for fruit jars, surgical instruments, cook- ing utensils, tubes, the framework and fittings of boats and air ships, telephone receivers, scientific apparatus, parts of opera glasses and telescopes, the framework of cameras, stock patterns for foundry work, and hardware samples. Its attractive appearance has led to its extensive use as an ornamental metal, both in interior decorative work and in numerous small objects, such as trays, picture frames, hairpins, and combs. Aluminium leaf is used for decorat- ing book covers and signs ; the powder is likewise used as a protective and attractive coating for letter boxes, steam pipes, lamp-posts, radiators, smokestacks, and other metal objects exposed to heat or the weather. During the last few years aluminium wire has come into use as a conductor of electricity. Large quantities of aluminium are used to reduce oxides, to make iron and steel more fluid, and to produce sounder castings. The applications of aluminium are constantly increasing. Alloys. The alloy of aluminium and copper aluminium bronze has been been described (see Alloys of Copper) . Magnalium is a recent alloy containing from 75 to 90 per cent of aluminium, the rest being magnesium. Aluminium Oxide, A1 2 O 3 , is the only oxide of alumin- ium. It is often called alumina, as silicon dioxide is called silica. Its native forms, corundum and emery, are found in Massachusetts, New Jersey, Georgia, Pennsylvania, North Carolina, and Canada ; large quantities come from Asia Minor and the islands near Greece. Emery is ex- tremely hard, and is used in various forms powder, cloth, paper, and wheels to grind and polish hard metals, plate Aluminium. 347 glass, etc. The crystallized varieties of aluminium oxide are usually known as corundum, and the transparent, colored kinds have long been prized as gems (see below). Alumina may be prepared by burning the metal or by heating its hydroxide. Thus prepared, it is a white powder, insoluble in water, but soluble in zfcids and in the caustic alkalies. It melts in the oxyhy- drogen flame, and in the electric furnace. Heating lessens its chemical activity. When alumina or any other compound of aluminium is heated, then cooled and moistened with cobaltous nitrate solution and heated again, the mass turns a beautiful blue color. This is a test for alu- minium. Aluminium is both basic and acid, that is, with acids it forms salts, like aluminium chloride, while with bases it forms aluminates. Gems containing Aluminium. Corundum (A1 2 O 3 ) has long been found as crystals in Ceylon, Siam, Burma, and other places in the Orient. The color is due to traces of impurities, usually oxides of metals. The sapphire is blue, and the ruby is red. The Oriental topaz is yellow, the Oriental amethyst is purple, and the Oriental emerald is green. Montana furnishes many sapphires, the output in 1901 being valued at $90,000. These gems may be artificially produced by dissolving alumina in a fused substance, adding an oxide to secure the desired color, and then allowing the alumina to crystallize. Spinels are complex compounds of aluminium. The typical or ruby spinel is magnesium aluminate (MgAl 2 O 4 ). It resembles the true ruby. Other spinels differ from the ruby spinel both in color and in composition. Turquoise is a complex aluminium phosphate containing traces of cop- per. It has a beautiful robinVegg-blue color, is compact, and may be worked into various shapes. Formerly turquoise came almost exclu- sively from Persia, but now New Mexico meets all demands. Nearly $120,000 worth of turquoise are mined annually m that state. Topaz is a complex aluminium silicate containing fluorine. It is usually pale yellow, and is found in many localities. Emerald is, next to diamond and ruby, the most precious gem. It is an aluminium silicate con- taining the rare element beryllium. The finest specimens have a deep emerald-green color and come from Colombia, South America. Garnet is a complex silicate of aluminium and another metal, especially cal- cium, magnesium, iron, or manganese. The kind used as a gem has a deep red color and is rather abundant. 348 Descriptive Chemistry. Aluminium Hydroxide, A1(OH) 3 , is a white, jelly like solid formed by adding an hydroxide to the solution of an aluminium salt, thus AlClg + 3 NH 4 OH = A1(OH) 3 + 3 NH 4 C1 Aluminium Ammonium Aluminium Ammonium Chloride Hydroxide Hydroxide 'Chloride It is insoluble in water. It interacts with strong acids and with alkalies (except ammonium hydroxide), forming respectively aluminium salts and aluminates. Thus A1(OH) 8 + 3 HC1 = A1C1 8 4- 3 H 2 O Aluminium Hydrochloric Aluminium Water Hydroxide Acid Chloride A1(OH) 3 + 3 NaOH = Na 3 AlO 3 + 3 H 2 O Sodium Sodium Hydroxide Aluminate Bauxite is a native aluminium hydroxide, though it contains iron and silicon. It resembles clay in texture and color. The vast deposits found at Baux, in southern France, furnish most of the raw material for the manufacture of aluminium, though about twenty thousand tons are annually obtained from our Southern states, chiefly from Georgia. Aluminium Sulphate, A1 2 (SO 4 ) 3 . 18 H 2 O, is a white, crystalline solid. The commercial salt has a variable com- position ; and, if pure, it dissolves readily and completely in water. It is extensively used in dyeing and paper making, and in preparing other aluminium compounds. Aluminium sulphate is prepared from pure clay, bauxite, or cryolite. If clay or bauxite is heated with sulphuric acid and then allowed to cool, the product is impure aluminium sulphate, known as " alum cake," or, if much iron is present, as " alumino ferric cake. 1 ' It is used to purify sewage and for other purposes where iron and the other impuri- ties do no harm. Purer aluminium sulphate is prepared by heating Aluminium. 349 bauxite with soda ash, extracting the sodium aluminate formed with water, and precipitating the aluminium, as the hydroxide with carbon dioxide gas. The relatively pure hydroxide is then changed into sul- phate by treatment with sulphuric acid. The product, known as "concentrated alum, 1 ' has the composition expressed by the formula A1.,(SO 4 ) 3 . 20 H 2 O, though separate crystals contain only eighteen molecules of water of crystallization. By boiling cryolite with milk of lime, the sodium aluminate thereby formed may be changed into " con- centrated alum," as described above. About 50,000 tons of "con- centrated alum " are annually produced in the United States. Alum. When solutions of aluminium sulphate and potas- sium sulphate are mixed and concentrated by evaporation, transparent, colorless, glassy crystals are deposited. This solid is potassium alum, or simply alum. It has the com- position represented by the formula, K 2 A1 2 (SO 4 ) 4 . 24 H 2 O, or K 2 SO 4 , A1 2 (SO 4 ) 3 . 24 H 2 O, and is sometimes called a double salt. It is the type of a class of similar salts called alums, which can be formed by crystallization from a mixture of aluminium sulphate and an alkaline sulphate. Alums are very soluble in water, and their solutions have an acid reaction and a sweetish, puckery taste. They crystallize alike, and contain twenty-four molecules of water of crystallization. When heated, alums lose their water of crystallization and some sulphuric acid, and fall to a white powder or porous mass known as burnt alum. Potassium alum is the most common, but ammonium and sodium alums are manufactured and used. Sodium alum is an ingredient of some baking powders. Burnt alum finds application as a medicine. Alum has been largely displaced by " concentrated alum," but the real alum is still used in dyeing and printing cloth, in tanning and paper making, in purifying water and sewage, as a medi- cine, for hardening plaster, in making wood and cloth fire- proof, and in preparing other aluminium compounds. 350 Descriptive Chemistry. Alum was known to the ancients, who used it in dyeing and tanning, and as a medicine. It was first manufactured in Europe, about the thirteenth century, from native alunite, which is an impure sulphate of aluminium, potassium, and iron. Alunite and alum slates or shales are now used to some extent, but most of the alum is made from bauxite. Not all alums contain aluminium. This metal may be replaced by iron, chromium, manganese, or similar metals, producing salhich have the same general properties as ordinary alum. formula of alums is M 2 (SO 4 ) 3 . X 2 SO 4 . 24 H 7 O, in aluminium, iron, chromium, etc., and X a metal (or group) like potas- sium, sodium, ammonium. Chrome alum (K.,Cr 2 (SO 4 ) 4 . 24 H.,0) belongs to this class. It is a purple, crystallized solid. The other alums have a limited, industrial application. * Alums and other aluminium salts are used as mordants in dyeing and calico printing. Some dyes must be fixed in the fabric by a metallic substance, otherwise the color would be easily removed. The cloth to be dyed or printed is impregnated or printed with the mordant, and then heated or treated with some substance to change the mor- dant into an insoluble compound. The mordanted cloth is next passed through a vat containing the solution of the dye, which unites chemically or mechanically (perhaps both) with the metallic compound, forming a colored com- pound. The latter is called a "lake"; it is relatively in- soluble, and cannot be easily washed from the cloth, i.e. it is a fast color. Aluminium acetate or "red liquor" and aluminium sulphate, besides alum, are used as mordants for cotton, linen, and wool. Cryolite is a white, glassy, crystallized solid. It often resembles clouded ice, and its name means "ice stone." Its composition corresponds to the formula Na 3 AlF 6 (or A1F 3 . 3 NaF). Small fragments melt easily, even in a candle flame, and color the Bunsen flame yellow. The only locality where it is found in commercial quantities is Aluminium. 351 southern Greenland, which yields annually about 10,000 tons. It is used not only in manufacturing aluminium, but as a source of alum and aluminium hydroxide, pure sodium carbonate and hydroxide, hydrofluoric acid, fluor- ides, and one kind of glass. Aluminium Chloride when pure is a white powder, but it is often a yellowish, crystalline mass (A1C1 3 . 6 H 2 O). It is prepared by heating powdered aluminium in chlorine, or by passing chlorine over a heated mixture of aluminium oxide and carbon. Exposed to the air, it absorbs moisture and gives off fumes of hydrochloric acid. It dissolves in water with evolution of heat, and if the solution is heated, hydrochloric acid is expelled, owing to the transformation of the chloride into the hydroxide, thus A1C1 3 + 3 H 2 = 3 HC1 + Al(OH), Aluminium Water Hydrochloric Aluminium Hy- Chloride Acid droxide This salt is used in organic chemistry. Clay is a more or less impure aluminium silicate, formed by the slow decomposition of rocks containing aluminium, especially feldspar. Pure feldspar is a silicate of alumin- ium and sodium or potassium. The products of its decom- position are chiefly an insoluble aluminium silicate and a soluble alkaline silicate. The latter is washed away. The aluminium silicate which remains is pure clay or kaolin. The latter is really a hydrous silicate, having the composi- tion corresponding to the formula Al 2 Si 3 O 7 , 2 H 2 O. The composition of clay varies, because it is seldom formed from pure feldspar. Most kaolin contains particles of mica and quartz. Ordinary clay contains many impurities, e.g. carbonates of calcium and magnesium, quartz, and iron compounds. Kaolin is a white, powdery mass. It becomes slightly plastic when wet, and can therefore be molded into various shapes. Ordinary clay is very plastic when Descriptive Chemistry. wet, more easily fused than kaolin, but shrinks consider- ably when dried and burned ; it also contains iron com- pounds, which color it gray, blue, yellow, brown, and red. All clays have a peculiar clayey odor when moist. Clay is the basis of pottery, of which there are three general kinds : porcelain or china, stoneware, and earthen- ware. Porcelain is the finest kind. It is made by heating to a high tem- perature a mixture of kaolin, fine sand, and some fusible substance, such as feldspar, chalk, or gypsum. The mass when cool is hard, dense, white, and translucent (if thin) ; it is not easily corroded by chemicals (ex- cept fused alkalies). Although it is not very porous, its surface is glazed, partly for protection, partly for ornament. This is done by coating it with a mixture similar to that used for making the porcelain but more easily fused, and then heating again so that the glaze will penetrate the surface. Stoneware is similar to porcelain, but coarser, because the materials are less carefully selected and prepared, and are not heated to such a high temperature. The best grades can hardly be distinguished from porcelain, but usually stoneware is much heavier and thicker. The cheaper kinds are made into jars, jugs, and bottles, especially large ones used in acid manufactories. Crockery is a fine grade of stoneware, though the best crockery is much like porcelain. If less pure, plastic clay is used and heated to a moderate temperature, the product is known as earthenware. This is a large class and in- cludes majolica, tiles, terra cotta, jugs, flowerpots, clay tobacco pipes, drain pipe, and bricks. This ware is porous and is usually glazed by throwing salt into the baking oven just before the operation is over. The salt volatilizes arid forms a fusible sodium aluminium silicate upon the surface. Cheap bricks are made from very impure clay, and their red color is due to iron oxides formed from the iron compounds in the unburned clay. Buff bricks are 'made from clay containing little or no iron, and clay containing silica yields fire-clay bricks, stove linings, retorts, and crucibles. EXERCISES. 1. What is the symbol and atomic weight of aluminium ? 2. Name several compounds of aluminium and discuss their occur- rence. What proportion of the earth's crust is aluminium ? Aluminium. 353 3. State briefly the history of aluminium. 4. Describe the metallurgy of aluminium by (#) the Hall process, () the Heroult process, (V) the older chemical method. 5. Discuss the production and cost of aluminium. 6. (#) Summarize the properties of aluminium. (<) State its uses. (V) Describe its alloys. 7. What is the formula and chemical name of alumina ? Describe its preparation. State its properties and uses. 8. State the properties and uses of corundum and emery. Review carborundum (see Chapter X). 9. Name seven gems containing aluminium. Describe them. 10. Describe aluminium hydroxide. How does it interact with acids and with alkalies ? 1 1 . What is bauxite ? For what is it used ? 12. Describe aluminium sulphate. State its properties and uses. How is it prepared ? What is " alum cake " ? u Alumino ferric cake " ? State their uses. 13. What is ordinary alum ? How is it manufactured ? State the general properties and uses of alums. What is (a) "concentrated alum, 1 ' and (^) burnt alum ? 14. Define a mordant. Describe its use. Name several mordants. What is (a) a " lake," (b) red liquor ? 15. What is the general formula of an alum ? What is chrome alum ? 16. Where is cryolite found ? State its properties and uses. What is its formula ? 17. Describe the preparation and state the properties of aluminium chloride. 18. What is clay ? How is it formed ? What is kaolin ? Describe (a) ordinary clay, and (6) kaolin. 19. Describe the manufacture of (a) porcelain, () stoneware, and (V) earthenware. Give an example of each. What is meant by glazing ? PROBLEMS. What is the per cent of aluminium in (a) cryolite (AlNa 3 F 6 ), () turquoise (A1,P 2 O 8 . H r Al 2 O 6 . 2 H 2 O), (V) corundum (A1 2 O 3 ), (W) aluminium hydroxide (A1(OH) 3 ) ? CHAPTER XXVI. TIN AND LEAD. TIN and lead are familiar metals. They have similar and useful properties, which give these metals and their compounds numerous applications. TIN. Occurrence of Tin. Metallic tin is rarely if ever found. Tin dioxide (cassiterite, tin stone, SnO 2 ) is the only available ore. It is not widely distributed, but large deposits are found in England (at Cornwall), Germany (in Bohemia and Saxony), Australia, Tasmania, and the East Indian Islands, especially Banca and Billiton. A small quantity is found, but not mined, in the United States. Tin is one of the oldest known metals. It is mentioned in the Pen- tateuch, and was obtained long before the Christian era by the Phoeni- cians from the British Isles, which were called Cassiterides (from the Greek word kassiteros, meaning tin). Many ancient bronzes contain tin. The alchemists called it Jupiter, and used the metal and its com- pounds. The Latin word stannum gives us the symbol Sn and the terms stannous and stannic. Metallurgy of Tin. If the tin ore contains sulphur or arsenic, these impurities must be removed by roasting. The tin oxide is then reduced by heating it with coal in a reverberatory furnace ; the simplest equation for this change is SnO 2 + C = Sn + CO 2 Tin Dioxide Carbon Tin Carbon Dioxide 354 Tin and Lead. 355 The molten tin which collects at the bottom of the furnace is drawn off and cast into bars or masses, which are often called block tin. Usually it is purified by melting it slowly on a hearth, inclined so that the more easily melted tin will flow down the hearth and leave the metallic impuri- ties behind. This tin may be further purified by stirring the molten metal with a wooden pole, or by holding billets of wood beneath its sur- face. The impurities which are oxidized by the escaping gases collect as a scum on the surface and are removed. Properties of Tin. Tin is a white, lustrous metal, which does not tarnish easily in the air. It is soft and malleable, and can be readily cut and hammered. It is softer than zinc, but harder than lead. Its specific gravity is 7.3. Tin may be obtained in the crystalline form, and when a piece of such tin is bent it makes a crackling sound, which is caused by the friction of these crystals upon one another. It melts at about 232 C, and when heated to a higher temperature' it burns, forming white tin oxide (SnO 2 ). The physical properties of tin, like those of zinc, vary with the temperature. Concentrated hydro- chloric acid changes it into stannous chloride (SnCl 2 ); treated with hot concentrated sulphuric acid, it forms stannous sulphate (SnSO 4 ) and sulphur dioxide ; and com- mercial nitric acid oxidizes it, the white, solid product being known as metastannic acid. Zinc precipitates tin from its solutions as a grayish black, spongy mass, which is sometimes filled with bright scales. Uses of Tin. Tin is so permanent in air, weak acids (like vinegar and fruit acids), and alkalies that it is exten- sively used as a protective coating for metals. Ordinary tinware is sheet iron coated with tin. The tin plate (sheet tin, or simply "tin") is made by dipping very clean sheet iron into molten tin. Tacks, nails, and many small iron objects are similarly tinned. Copper coated with tin 356 Descriptive Chemistry. is made into vessels for cooking, and brass coated with tin is made into pins. Large quantities of tin plate are used to cover roofs. Tinned iron does not rust until the tin is worn off and the iron exposed, and then the rusting proceeds rapidly. Tin is also hammered into thin sheets called tin foil, though much of the tin foil now used con- tains lead. Many useful alloys contain tin as an essential ingredient. During the last few years the annual con- sumption of tin has been about 75,000 pounds. Alloys of tin are described under COPPER. Those containing a minor percentage of tin are .bronze, gun metal, bell metal, speculum metal, type metal, anti-friction metals, and fusible alloys. Britannia metal contains about 90 per cent tin, 8 per cent antimony, and the rest mainly copper. It is a white metal, and was formerly made into tableware. White metal contains less tin and more antimony than Britannia, though the composition varies. It resembles Britannia. The harder varieties of 'white metal are used as parts of machinery, and the softer kinds are made into ornaments and cheap jewelry. Pew- ter and solder contain varying proportions of tin and lead. Plumbers' solder, or soft solder, is about one third tin and two thirds lead. It is harder than either constituent, but it melts at a lower temperature. Tin amalgam is sometimes used to coat mirrors. Compounds of Tin. Tin forms two series of compounds, the stan- nous and the stannic. Stannic oxide (SnO 2 ) has already been men- tioned as the chief ore of tin, and as the product formed when tin is burned. The artificial oxide is faint yellow when hot and white when- cold. The native oxide is a brown or black, lustrous, and often crystal- lized solid. Irregular pebbles called stream tin occur in some localities near rivers. Stannous chloride (SnCl.,) is formed by the interaction of hydrochloric acid and tin. From the concentrated solution a green- ish salt crystallizes (SnCl 2 . i H.,O), known as tin crystals or salt of tin. Tin and Lead. 357 Stannous chloride passes readily into stannic chloride (SnCl 4 ) when added to mercuric chloride solution. The simplest equation for this change is SnCl 2 + 2 HgCl 2 = SnCl 4 + Hg 2 Cl 2 Stannous Mercuric Stannic Mercurous Chloride Chloride Chloride Chloride By an extension of the simplest idea of oxidation and reduction, the stannous chloride in the change is said to be oxidized to stannic chlo- ride, but it reduced the mercuric chloride to mercurous chloride. Stan- nous chloride is often used as a reducing agent and as a mordant in dyeing and calico printing. Crystallized stannic chloride (SnCl 4 . 5 H 2 O), known commercially as oxymuriate of tin, is also used as a mordant. Tin mordants produce brilliant colors. Sodium stannate (Na 2 SnO 3 . 3 H 2 O) is extensively used to prepare cotton cloth for printing. LEAD. Occurrence of Lead. Metallic lead is occasionally found in small quantities. The most abundant ore is lead sulphide (galena, PbS). Other native compounds, formed by the alteration of galena, are the carbonate (cerussite, PbCO 3 ), the sulphate (anglesite, PbSO 4 ), and the phos- phate (pyromorphite, Pb 5 Cl(PO 4 ) 3 ). Lead compounds are widely distributed, but the source of commercial lead is the sulphide. Lead has been used by civilized people since the dawn of history. The Chinese have used it for ages to line chests in which tea is stored and transported. The Romans, who obtained it from Spain, called it plumbum nigrum, i.e. black lead. The symbol Pb Qomes from plumbum. The ancients also used lead compounds (especially the carbonate and red oxide) as paints and cosmetics. The annual production of lead has increased rapidly during the last few years, and in 1902 it was about 800,000 tons. This vast amount comes chiefly from the United States, Spain, Germany, Mexico, New South Wales, and England. The United States in 1902 produced about 250,000 tons of lead from ores found mainly in the Middle West (Illinois, Iowa, Wisconsin, and Missouri), Colorado. Idaho, and Utah. 358 Descriptive Chemistry. Metallurgy of Lead. Lead is readily obtained from galena, (i) In the reduction process the ore is roasted in a reverberatory furnace until a part of the sulphide is changed into lead oxide and lead sulphate. The equations for these changes are 2 PbS 4- 3 O 2 = 2 PbO + 2 SO 2 Lead Sulphide Oxygen Lead Oxide Sulphur Dioxide PbS + 2O 2 PbS0 4 Lead Sulphide Oxygen Lead Sulphate The air is then shut off and the mixture of the three lead compounds is heated to a higher temperature. By this operation the lead sulphide interacts with the other lead compounds, forming lead and sulphur diox- ide, thus 2 PbS + PbSO 4 + 2 PbO = sPb + 3 SO 2 Lead Sulphide Lead Sulphate Lead Oxide Lead Sulphur Dioxide (2) Ores poor in lead are sometimes reduced by roasting with iron, which combines with the sulphur, leaving the lead free, thus PbS + Fe = Pb + FeS Lead Sulphide Iron Lead Iron Sulphide (3) At Niagara Falls lead is obtained from galena by electrolysis. Crushed galena is made the cathode, dilute sulphuric acid is the electro- lyte, and the bottom of the reduction pan is the anode. The sulphur is changed into hydrogen sulphide, which escapes into a combustion chamber where its sulphur is recovered or converted into sulphuric acid. The lead remains in the pan as a spongy mass. The silver, which remains in the lead obtained by reduction, is extracted by the Parkes process (see Silver). Properties of Lead. Lead is a bluish metal. When scraped or cut, it has a brilliant luster, which soon disap- pears, owing to the formation of a film of oxide. This coating protects the lead from further change. It is a soft metal, and may be scratched with the finger nail. It dis- colors the hands, and when drawn across a rough surface it leaves a black mark. For this reason it is sometimes Tin and Lead. 359 called black lead (see Graphite). Lead is not tough enough to be readily hammered into foil or drawn into fine wire, but it can be rolled into sheets. It is a heavy metal, its specific gravity being 11.35; w ith the exception of mercury, it is the heaviest of the familiar metals. It melts at 326 C, or about 100 higher than tin and 100 lower than zinc. Lead, when heated strongly in air, changes into an oxide (mainly the monoxide, PbO). Hydrochloric and sulphuric acids have little effect upon compact lead. Nitric acid changes it into lead nitrate (Pb(NO 3 ) 2 ). Acetic acid (or vinegar) and acids from fruits and vegetables change it into soluble, poisonous compounds ; hence cheap tin-plated vessels, which sometimes contain lead, should never be used in cooking. Zinc and iron precipitate lead from its solutions as a grayish mass, which often has a beautiful treelike appearance. Lead in Drinking Water. Lead is slowly changed into soluble compounds by water containing carbon dioxide, ammonia, nitrates, or chlorides. But water containing sul- phates or carbonates forms an insoluble coating on the lead, thus protecting it from further action. All lead salts are poisonous, and if taken into the system they will slowly accumulate and ultimately cause serious and dangerous illness. Water suspected of attacking lead should never be drunk after it has been standing very long in lead pipes, but should be allowed to flow until the pipe has been filled with fresh water. Sometimes the water cannot be drunk at all. The city of Lowell, Massachusetts, recently aban- doned one source of its water supply because of the rapid solvent action of the water upon lead pipes. Uses of Lead. Lead is extensively used as pipe, be- cause it can be made into indefinitely long pieces, which Descriptive Chemistry. can be easily bent, cut, and united (by solder). The pipe is made by forcing softened lead through a hole in a steel plate or by the apparatus shown in Figure 69. Lead pipe is not only used to convey water to and from parts of build- ings, but as a sheath for copper wires, both overhead and underground. As sheet lead it is used to cover roofs and to line sinks, cisterns, and the cells employed in many electrolytic processes. The lead chambers and evaporating pans used in manufacturing sulphuric acid are made of sheet lead. Shot and bullets are lead (alloyed with a little arsenic). Spongy lead is used in preparing inthelongcylin- , r der.cc, is forced tne plates of storage batteries. K'tough The A11 y s of Lead are important. Type metal contains 70 to 80 per cent lead ; the FIG. 69. Ap- Ing* lead' pipe! The molten lead the space, D, varied insL by other constituents are tin and antimony. The the steel rod, A. latter metal expands when it solidifies and makes the face of the type sharp and clear. Solder, pewter, and fusible alloys contain lead as an essential constituent (see Alloys of Tin). Small quantities are found in brass and bronze. Lead Oxides. There are three important oxides. Lead monoxide (PbO) is a yellowish powder known as massicot, or a buff-colored crystalline mass called litharge. It is formed by heating lead above its melting point in a cur- rent of air. It is made this way, though considerable is obtained as a by-product in separating silver from lead (see Cupellation). Large quantities are used in preparing some oils and varnishes, flint glass, other lead compounds, and as a glaze. Lead tetroxide (red lead, minium, Pb 3 O 4 ) is a red powder, 'varying somewhat in color and Tin and Lead. 361 composition. It is prepared by heating lead (or lead mo- noxide) to about 350 C. It is used in making flint glass. Pure grades are made into artists' paint, but the cheap variety is used to paint structural iron work (bridges, gasometers, etc.), hulls of vessels, and agricultural imple- ments. It is used in plumbing and gas fitting to make joints tight. Orange mineral has the same composition as red lead, and although its color is lighter, its uses are the same. Lead dioxide (lead peroxide, PbO 2 ), is a brown powder formed by treating lead tetroxide with nitric acid. It is used in storage batteries. Lead Carbonate, PbCO 3 , is found native as the trans- parent, crystallized mineral cerussite. It is obtained as a white powder by adding ammonium carbonate solution to lead nitrate solution. Sodium and potassium carbonates, however, form basic lead carbonates, which have a compo- sition depending upon the temperature. The most im- portant of these basic carbonates has the composition corresponding to the formula 2 PbCO 3 . Pb(OH) 2 , and is known as white lead. It is a heavy, white powder which mixes well with linseed oil, and is used extensively as a white paint and as the basis of many colored paints. White lead is manufactured by several processes. The Dutch process is the oldest, having been used as early as 1622. It is essentially the same to-day, though many details have been improved. Perforated disks of lead are put in earthenware pots which have a separate com- partment at the bottom, containing a weak solution of acetic acid (about as strong as vinegar). These pots are arranged in tiers in a large brick building, and spent tan bark is placed between each tier. The building is now closed except openings for the entrance and exit of air and steam. The heat volatilizes the acetic acid which changes the lead into a lead acetate. The tan bark ferments and liberates car- bon dioxide, which changes the lead acetate into basic lead carbonate or white lead. The whole operation requires from sixty to one hun- dred days. The slowness is the chief objection to this process. In 362 Descriptive Chemistry. the German process acetic acid vapor, steam, and carbon dioxide are forced into closed chambers in which sheets of lead are suspended. It requires about five weeks. In the French process basic lead carbonate is precipitated from a basic lead acetate by carbon dioxide. Milner's process is a modification of the French process. Both are quicker than the Dutch or German processes, but the product is not considered so good. An electrolytic process has recently been devised. The anode is lead, the cathode is copper, and the electrolyte is sodium nitrate solution. When the electric current is passed, (i) nitric acid is liber- ated at the anode, and changes the lead into lead nitrate, and (2) at the cathode sodium is formed, which decomposes the water, thereby forming sodium hydroxide. The lead nitrate and sodium hydroxide solutions interact, forming insoluble lead hydroxide and sodium nitrate, thus Pb(NO 3 ) 2 + 2NaOH = Pb(OH) 2 + 2 NaNO 3 Lead Nitrate Sodium Hydroxide Lead Hydroxide Sodium Nitrate The sodium nitrate is left in the cell to be acted upon again, but the lead hydroxide is changed into lead carbonate by treatment with sodium bicarbonate. This process is rapid, and the product is claimed to be as good as white lead produced by other processes. White lead paint often turns dark in the air, owing to the formation of lead sulphide, which is black. Its extensive use is largely due to its great covering power, i.e. a very thin layer produces a perfectly white surface, and therefore less paint is required for a given area. It is often adulterated with zinc oxide and barium sulphate; those are white solids, but they are cheaper and have less covering power. Lead Sulphide, PbS. Native lead sulphide is the min- eral galena, the chief ore of lead. It resembles lead in FIG. 70. Galena crystals (cube, octahedron and cube, octahedron). appearance, but is harder and is usually crystallized as cubes, octahedrons, or their combinations (Fig. 70). It Tin and Lead. 363 has perfect cubic cleavage, i.e. it breaks into cubes or frag- ments more or less rectangular. It is easily changed into lead by heating it alone or with sodium carbonate on char- coal. Lead sulphide, as prepared in the laboratory, is a black solid. Black lead sulphide is readily precipitated from a lead salt solution by hydrogen sulphide. Its formation is the test for lead. It is changed into lead chloride by concentrated hydrochloric acid and into lead sul- phate by concentrated nitric acid. Other Compounds of Lead, which are important, are the chloride, sulphate, nitrate, chromate, and acetate. Lead chloride (PbCl 2 ) is a white solid formed by adding hydrochloric acid or a soluble chloride to a cold solution of a lead salt. It dissolves in hot water. Lead sul- phate (PbSO 4 ) is a white solid, formed by adding sulphuric acid or a soluble sulphate to a solution of a lead salt. It is very slightly soluble in water, but soluble in concentrated sulphuric acid, hence crude sul- phuric acid often contains lead sulphate. Lead nitrate (Pb(NO 3 ) 2 ) is a white crystallized solid formed by dissolving lead (or better, lead mo- noxide) in nitric acid. When heated, it decomposes into lead oxide (PbO), nitrogen peroxide, and oxygen. Lead acetate (Pb(C 2 H 3 O 2 ) 2 ) is a white, crystallized solid formed by the action of acetic acid upon lead or lead oxide (PbO) . It is very soluble in water and is often called " sugar of lead. 1 ' EXERCISES. 1. Name the chief ore of tin. Where is it found? What is " stream tin"? 2. Give briefly the history of tin. What is its symbol ? Why? 3. Describe (a} the metallurgy of tin, and (6) its purification. 4. Summarize the properties of tin. State its* uses. 5. What is "tin 11 ? Block tin? Tinfoil? Tinware? Sheet tin? Tin plate ? 6. Describe three alloys which contain large proportions of tin. Name several alloys containing a minor proportion of tin. 7. Compare native and artificial tin oxide (SnO 2 ). 8. What is the formula of (a} stannous chloride, and (b) stannic chloride? What is their chemical relation? State the use of each chloride. What other names has stannous chloride? 364 Descriptive Chemistry. 9. What is the most abundant ore of lead? Name other native compounds. 10. Give a brief history of lead. What is its symbol? Why? 11. Discuss the production of lead. 12. Describe the metallurgy of lead by (a) the reduction process, () roasting with iron, (c) electrolysis of galena. 13. Summarize the properties of lead. 14. State the uses of lead. 15. Discuss the relation of lead to water. 1 6. What is (a) type metal, (6) solder, (c) fusible alloy? 17. Give the name and formula of the oxides of lead. 1 8. Describe the preparation, and state the properties and uses of (a) litharge, (#) red lead, (c) lead peroxide. 19. What is white lead? Describe its preparation by (a) the Dutch method, and () electrolysis of sodium nitrate. 20. State the properties and uses of white lead. 21. What is the formula and chemical name of galena? Describe this mineral. Describe the corresponding artificial compound. What is the test for lead? 22. Describe the following salts of lead : (a) chloride, (b) sulphate, (c) nitrate, (d) acetate. PROBLEMS. 1. What is the per cent of lead in (a} galena (PbS), () cerussite (PbCO 8 ), (c) anglesite (PbSO 4 ), (d) lead acetate (Pb(C 2 H 3 O 2 ) 2 . 3 H 2 O) ? 2. How much litharge may be made from 40.5 gm. of lead? (As- sume Pb + O = PbO.) 3. What is the per cent of tin in (a) tinstone (SnO 2 ), (b) stannous chloride (SnCl 2 ), (c) stannic chloride (SnCl 4 )? CHAPTER XXVII. CHROMIUM AND MANGANESE. THESE elements do not belong to the same group, but they have several common properties and form analogous compounds. CHROMIUM. Occurrence of Chromium. Metallic chromium is never found free. Its chief ore is an oxide (chromite, chrome iron ore, FeCr 2 O 4 ). Native lead chromate (crocoite or crocoisite, PbCrO 4 ) is less common. Traces of chromium occur in many green minerals and rocks, e.g. emerald and serpentine, and verde antique marble. Chromite is mined chiefly in Greece, New Caledonia, New South Wales, Turkey, and Canada. The total annual production is about 30,000 tons. The word chromium comes from the Greek word chroma, meaning color, and emphasizes the fact that most chromium compounds have decided colors. Preparation, Properties, and Uses. Chromium was a rare metal until Moissan prepared it, in 1894, in the electric furnace. Now it is produced in quantities by heating a mixture of chromite and carbon in an electric furnace. The crude chromium is refined by fusing it with lime. Very pure chromium is also prepared by reducing chromic oxide with aluminium powder. Chromium is a lustrous gray metal. It takes a good polish, which is not removed by exposure to air. It is hard, but it can be filed and pol- ished without difficulty. Its specific gravity is about 6.9. It is not attracted by a magnet. It can be fused only in the electric furnace. Chromium is used to harden the steel, which is to be made into armor, projectiles, safes, and vaults, and parts of machines used to 365 366 Descriptive Chemistry. crush gold-bearing quartz. This hardened steel is called chrome steel. The commercial form of chromium is an alloy of 65 to 80 per cert chromium, a little carbon, and the rest iron ; this alloy is called ferro- chrome. Compounds of Chromium are numerous, some are com- plex, many pass readily into one another, and a few have industrial applications. The most important are potassium chromate, potassium dichromate, chrome alum, and lead chromate. Potassium Chromate (K 2 CrO 4 ) and Potassium Dichro- mate (or Bichromate, K 2 Cr 2 O 7 ). These compounds are manufactured from chrome iron ore. The crushed ore is mixed with lime and potassium carbonate, and roasted in a reverberatory furnace ; air is freely admitted and the mass is frequently raked. By this operation the ore is oxidized into a mixture of calcium and potassium chro- mates. The mass is cooled, pulverized, and treated with a hot solution of potassium sulphate, which changes the calcium chromate into potassium chromate. The clear, saturated solution of potassium chromate is changed by sulphuric acid into potassium dichromate ; the latter is purified by recrystallization from water. Potassium chro- mate is a lemon-yellow, crystallized solid, very soluble in water. Acids change it into the dichromate, thus 2 K 2 CrO 4 + H 2 SO 4 = K 2 Cr 2 O 7 + K 2 SO 4 + H 2 O Potassium Sulphuric Potassium Potassium Water Chromate Acid Dichromate Sulphate Potassium Dichromate is a red solid which forms large crystals. It is less soluble in water than potassium chro- mate. Alkalies change it into a chromate, thus K 2 Cr 2 O 7 + 2KOH = 2 K 2 CrO 4 + H 2 O Potassium Potassium Potassium Water Dichromate Hydroxide Chromate Chromium and Manganese. 367 Potassium dichromate is used in dyeing, calico printing, and tanning, in bleaching oils, and in manufacturing other chromium compounds and dyestuffs. Its uses depend mainly upon the fact that it is an oxidizing agent. When hydrochloric acid is added to potassium dichromate, oxy-r gen from the dichromate withdraws hydrogen from the acid and liberates free chlorine, thus K 2 Cr 2 O 7 -f 14 HC1 = 2 KC1 + 2 CrCl 3 + 3 C1 2 + 7 H 2 O Potassium Di- Hydrochloric Potassium Chromic Chlorine Water chromate Acid Chloride Chloride If an oxidizable substance is present, such as organic mat- ter, alcohol, or a ferrous compound, it is quickly oxidized. Potassium chromate is also formed as a yellow mass by fusing on porcelain or platinum a mixture of a chromium compound, potassium carbonate, and potassium nitrate. When the mass is boiled with acetic acid to decompose the carbonate and expel carbon dioxide, and then added to a lead salt solution, yellow lead chromate is formed. This experiment is often used as a test for chromium. Chrome Alum, K 2 Cr 2 (SO 4 ) 4 . 24 H 2 O, is a purple, crys- tallized solid. It is analogous in composition and similar in properties to ordinary alum, but it contains chromium instead of aluminium. It can be prepared by mixing potassium and chromium sulphates in the proper propor- tion, or by passing sulphur dioxide into a solution of potassium dichromate containing sulphuric acid. The commercial substance is a by-product obtained in the manufacture of alizarine, a dye which yields magnificent colors. Chrome alum is used as a mordant in dyeing and calico printing, and in tanning. Lead Chromate, PbCrO 4 , is a bright yellow solid, formed by adding potassium chromate or dichromate to a solution of lead salt: It is known as chrome yellow and is used as the basis of yellow paint When boiled with sodium 370 Descriptive Chemistry. called black oxide of manganese. When heated it yields oxygen ; and when heated with hydrochloric acid the two compounds interact, forming manganous chloride, chlorine, and water, thus MnO 2 + 4HC1 = MnCl 2 + Cl a + H 2 O Manganese Hydrochloric Manganese Chlorine Water Dioxide Acid Chloride It colors glass and borax a beautiful amethyst, and" is often added to common glass to neutralize the green color. Enormous quantities are used in the manufacture of oxy- gen, chlorine, glass, and manganese alloys and compounds. The manganese dioxide used in the manufacture of chlorine is recov- ered by the Weldon process. The impure manganous chloride solu- tion from the chlorine still is treated with calcium carbonate to neutralize free acid and precipitate any iron present. Lime is added to the clear solution of manganous chloride, and air is blown into the mixture. The manganous chloride is changed into manganous hydroxide (Mn(OH).,), which interacts with the oxygen (of the air) and lime, forming chiefly calcium manganite (CaMnO 3 , or CaO . MnO 2 ). After this mixture has settled, the calcium chloride is drawn off, and the manganese compound, which is called " Weldon mud," is used to generate more chlorine. Manganese dioxide was used by the ancients to decolorize glass, but its nature was misunderstood. They confused it with an iron oxide called magnesia stone, and the alchemists in the Middle Ages gave the name magnesia to this manganese dioxide. Later they called it magnesia nigra, or black magnesia, to distinguish it from magnesia alba, or white magnesia (MgO), supposing that the two were related. Man- ganese was isolated in 1774, and later was given the specific name manganesium, which was soon shortened to manganese. Potassium Permanganate, KMnO 4 , is a dark purple, glistening, crystallized solid, though the crystals sometimes appear black, with a greenish luster. It is very soluble in water, and the solution is red, purple, or black, according to the concentration. Potassium permanganate gives up its oxygen readily and is used as an oxidizing agent in the Chromium and Manganese. 371 laboratory and on a large scale to purify stagnant water and sewage. It is such a powerful oxidizing agent that it cannot be filtered through paper, but only through asbestos or spun glass. It is also used as a disinfectant, as a medi- cine, in bleaching and dyeing, in coloring wood brown, and in purifying gases, such as hydrogen, ammonia, and carbon dioxide. Potassium permanganate is manufactured by oxidizing a mixture of manganese dioxide and potassium hydroxide, and treating the resulting potassium manganate with sulphuric acid, carbon dioxide, or chlorine. The essential reactions are represented thus MnO, + 2KOH + O = K,MnO 4 + H 2 O Manganese Potassium Potassium Dioxide Hydroxide Manganate 3 K 2 MnO 4 + 2 CO 2 = 2 KMnO 4 + K 2 CO 3 + MnO 2 Potassium Permanganate The uses of potassium permanganate depend mainly upon its oxidiz- ing power. With sulphuric acid the action is represented thus 2KMnO 4 + 3H 2 SO 4 = 50 + 2 MnSO 4 + K 2 SO 4 + 3 H 2 O Potassium Sulphuric Oxygen Manganese Potassium Water Permanganate Acid Sulphate Sulphate The liberated oxygen attacks at once any organic matter present, and the solution becomes brown or colorless, owing to the decomposition of the potassium permanganate into colorless compounds. Compounds of Manganese, like those of chromium, are numerous, often complex, and closely related. There are four oxides besides manganese dioxide. Three manganous compounds are important, the chloride (MnCL,), the sulphate (MnSO 4 ), and the sulphide (MnS). The chloride and sulphate are pink, crystallized salts, and the sulphide is a flesh-colored precipitate formed by adding ammonium sulphide to the solution of a manganous salt, thus distinguishing it from all other sulphides. Manganates are salts of the hypothetical manganic acid (H 2 MnO 4 ). They are analogous to chromates, and the manganese in them acts as a non-metal. Potassium manganate is obtained as a green mass by fusing a mixture of a manganese compound, potassium 372 Descriptive Chemistry. hydroxide (or carbonate), and potassium nitrate. Its formation on a small scale constitutes the test for manganese. Sodium manganate is used in solution as a disinfectant. EXERCISES. i . What is the symbol of chromium and of manganese ? Why is each element so named? 2. What is the chief ore of chromium? Where is it found? What other minerals contain chromium? 3. Describe the preparation of chromium. State its properties and uses. What is chrome steel? Ferrochrome? 4. Describe the manufacture of (a) potassium chromate, and () po- tassium dichromate. State their properties and uses. What is the formula of each? 5. What are the tests for chromium? 6. Describe chrome alum. How is it made? State its uses. How does it differ from ordinary alum? 7. Describe lead chromate. How is it formed ? For what is it used ? 8. In what two ways does chromium act in its compounds? What is chromic oxide? For what is it used? What is chromium trioxide? How is it related to potassium dichromate? 9. Name several ores of manganese. What is the chief ore ? Dis- cuss the production of manganese ores. 10. Describe the preparation, and state the properties of manganese. 11. What is spiegel iron? Ferromanganese? State their uses. 12. Describe manganese dioxide. State its properties and uses. How is it recovered by the Weldon process? What is the common name of manganese dioxide? Why is it so called? PROBLEMS. 1. What is the per cent of chromium in (a) lead chromate (PbCrO 4 ), () chrome ironstone (Cr 2 O 3 . FeO), (c) chromic oxide (Cr 2 O 3 ) ? 2. What is the per cent of manganese in ( K = 39. ^~ = 2 3 - The same is true of phosphorus, arsenic, and antimony P= 3 i, As= 75> Sb=i20. 3I + I2 = 7S . 5 . 398 Descriptive Chemistry. The existence of other relations similar to these, together with a deep desire to obtain more accurate atomic weights and a growing interest in the properties of the elements themselves, focused the attention of chemists at this time (1855-1865) upon the relation of properties to atomic weights. Several things fostered the above principle. One was the atomic weight determinations of Stas, whose masterly work proved beyond doubt that Prout was incor- rect when he insisted in 1815 that the atomic weights are whole numbers. Another was the acceptance by most chemists of the same table of atomic weights. A third was the rapid accumulation of many facts about the ele- ments and their compounds. Chemists were ready for a new classification of the elements. The Periodic Classification. Previous to 1869 no classification included all the elements. In that year the Russian chemist Mendeleeff published a classification of the elements according to the periodic law. His views had been partially anticipated by several chemists, and were soon amplified by the German chemist, Lothar Meyer. Their classification of the elements revealed a new relation between the properties of the elements and their atomic weights. If all the elements are arranged in the order of their increasing atomic weights beginning with lithium, their properties will vary periodically, i.e. at certain regu- lar intervals or periods elements will be found which have similar properties. In other words, a certain increase in atomic weight causes a reappearance or return of prop- erties. The general relation is often summarized in the Periodic Law - The properties of the elements are periodic functions of their atomic weights. General Relations of the Elements. 399 cu o , in ' ^ " co vo 1 w ro ^ iT IT !T if j i | ill 0. o O in ? CO g 2 II - II 8 II II 1 1 "1 1 ~ 1 1 1 1 O ^ ! ^ 2 1 S 1 1 ~ 1 I Q cu I O . . o? 4- H in N M ro f* M w> S II II II ^ II 1 II ^ ^ IT < r ^ ii ; s > ' 6 H 1 8 10 ON vd 5- f 1 | 1 l ' ' . * ? H N U 1 H O C4 t^. M O* M f ^ f -5 f ' | P ' co > i3 W 1 CU D O O v? S ' a .* . *;-. 1.- S 4-..I i i , 1, | N W o s d ^ p I i i o s hx ts. ^COOs^^ilcol (if iTnir 3 7 W) ii' 3 a i M ^ s * ^ L_i_J fi H M ^1 Or^QO Oi OH 400 Descriptive Chemistry. Function here means the exhibition of some special rela- tion, viz. that of properties to atomic weight. Interpreted freely, the law means (i) properties and atomic weight are related, they depend upon each other; and (2) this relation is exhibited again and again as we reach elements with increasing atomic weights at regular intervals in the suc- cessive arrangement The Periodic Table originally proposed by Mendeleeff has been modified from time to time, as new facts have necessitated. The table generally accepted at the present time is given on page 399. From the table it is seen that the elements fall naturally into two subdivisions, (i) Those in the same vertical col- umn belong to the same natural group or family. Thus, in Group I are found the alkali metals, in Group II the alkaline earth metals, in Group VII the halogens. (2) The elements in the same horizontal row belong to the same period. The periodic variation of their properties is well illustrated by the second and third periods. Begin- ning with lithium, the general chemical properties vary regularly with increasing atomic weight Thus, the metal- lic character gradually diminishes until fluorine is passed and sodium is reached; here it reappears. Proceeding onward from sodium, the same gradation of properties is noticed until potassium is reached, and here again the marked metallic character in the same way reappears. There is no sudden change in properties until we pass from one period to the next. Thus, fluorine at the end of the second period forms a powerful acid, but sodium at the beginning of the third period forms a strong base. Simi- larly, chlorine is strongly acidic ; but potassium, which begins the next period, is markedly basic; chlorine is a typical non-metal, while potassium is a typical metal. Not General Relations of the Elements. 401 all elements fit the periodic classification equally well, but the arrangement is at least very suggestive, and doubtless expresses an approximately truthful relation. The Gaps in the Periodic Classification probably corre- spond to elements not yet discovered. Three such gaps, which were in the original table, have been filled. When Mendeleeff proposed his arrangement, he predicted the discovery of three elements having definite properties. These elements, gallium, scandium, and germanium, have since been discovered and now occupy their pre- dicted place in the table. Possibly other gaps will be filled by newly discovered elements. The discovery of the predicted elements was not the only immediate service of MendeleefFs table. It also emphasized the necessity of more accurate atomic weights. Several elements did not fall into their proper places, and careful investigation showed that their accepted atomic weights were incorrect. Thus, the atomic weights of beryllium and in- dium were changed to their present values, and the present order of the platinum metals was adopted ; cobalt and nickel are still being studied. The position of argon, helium, and very rare metals is still doubtful, owing to a limited knowledge of their properties and atomic weights. Hydrogen, also, still lacks a place. SPECTRUM ANALYSIS. Introduction. When light from an ordinary gas flame, glowing lime or other solid, or a Welsbach flame is passed through a prism and falls upon a white surface, a long band of color is produced. The colors are perfectly blended, and are arranged like the familiar colors of the rainbow. This band of colors is called a spectrum. The white light has been separated or analyzed into the col- ored. The examination and study of the spectrum of a substance is spectrum analysis, and it is accomplished by a spectroscope. 402 Descriptive Chemistry. The Spectroscope consists essentially of a prism and tubes, one of which is a telescope (Fig. 75) . The light enters a slit in the tube, passes FIG. 75. A spectroscope. through, and falls upon the prism. Here it is bent from its path, and as it emerges from the prism, it may be viewed through the telescope as a magnified spectrum. Kinds of Spectra. (i) The spectrum of an incan- descent solid is a continuous band of colors. (2) But the spectra of gases are narrow, colored, vertical bars or lines, separated by black spaces. Thus, sodium vapor has a yel- low line, potassium a red and a violet line, and barium sev- eral lines where the green and yellow parts of the ordinary spectrum occur. Each element which is a gas, or which can be vaporized, has its own bright line spectrum. The lines always occupy the same relative positions, which in most cases have been very carefully determined. Therefore, when examined through a spectroscope, the yellow line of sodium will always be seen in its proper place, and the red and violet potassium lines in their places. Therefore, by examining the light from different substances, it is possi- General Relations of the Elements. 403 ble to tell what elements they contain. (3) The spectrum of sunlight is the familiar band of colors, but it is crossed vertically by many black lines, which have fixed positions (Fig. 76). It is believed that the sun is a glowing hot solid, surrounded by very hot gases. It therefore should AaBC D Eb F G If 111 ill I Hed Orange Fellow Green Blue Indigo Violet FIG. 76. Spectrum of sunlight showing some of the vertical lines. give the two kinds of spectra, the continuous and the bright line. Now it has been proved that the vapor of an element absorbs the light given out by the same ele- ment when solid. Hence the dark lines which appear in the solar spectrum are caused by the absorptive power of the gases in the sun's atmosphere. The solar spectrum is often called an absorptive spectrum. Spectrum Analysis. In the laboratory the spectro- scope is used to detect the presence of certain elements, more especially the metals. If the metal or one of its compounds is put on a platinum wire and held in the Bunsen flame before the slit, the characteristic spectrum of the element can be easily recognized in the telescope. Two spectra do not interfere, because each line has its own place. Hence several elements may be distinguished in a mixture. Minute quantities are easily detected by the spectroscope. Rare elements, which can be obtained only in very small quantities or with great difficulty, are studied by the spectroscope. Thus, Bunsen, who (with Kirch- hoff) devised the improved spectroscope, discovered the rare metals, rubidium and caesium. And within the last few years the spectroscope has been especially serviceable 404 Descriptive Chemistry. in studying argon, helium, krypton, neon, and xenon. By means of the spectroscope it has been shown that the sun contains many elements found in our earth. Accord- ing to a reliable authority, about thirty of the elements known to us are present in the sun. The spectroscope also enables astronomers to tell the nature of stars, comets, nebulae, and other heavenly bodies. The stars thus far examined give spectra crossed by dark lines, and therefore these bodies are like the sun ; but nebulae give bright line spectra, and hence consist of incandescent gases. EXERCISES. 1. Discuss the classification of the elements according to (a) metals and non-metals, () acid and basic properties, (c) valence, (d) groups based on resemblances, (e) numerical relations. 2. What is the fundamental idea of the periodic classification ? How does it differ from previous systems ? When and by whom was this classification proposed and developed ? 3. State the periodic law. Explain it. What is meant by (a) func- tion, () period, (c} group ? 4. Illustrate the law by (a) the alkali metals, and (b} the halogens. 5. Discuss the gaps in the periodic arrangement of the elements. 6. Of what use has this law been ? 7. State some objections to it. 8. Describe (a) a continuous spectrum, () a line spectrum, (c) an absorption spectrum. 9. Describe a spectroscope. How is it used ? 10. What kind of a spectrum is produced by (a} a glowing solid, () a glowing vapor, (c) a glowing solid surrounded by a glowing vapor? n. What is spectrum analysis ? How is it applied (a) in the labo- ratory, and (b) by astronomers ? 12. What does spectrum analysis show about each element ? About their relations to each other ? About their distribution ? About the heavenly bodies ? 13. Who perfected the spectroscope and developed its use ? 14. What recent use has been made of the spectroscope in (a) chem- istry, and (b) astronomy ? CHAPTER XXXI. SOME COMMON ORGANIC COMPOUNDS. Introduction. In the early days of chemistry it was believed that starch, sugar, and other compounds obtained from plants and animals were produced by the influence of some mysterious vital force. Such compounds were called organic, because of their connection with living things, i.e. with bodies having organs ; and they were sharply dis- tinguished from inorganic or mineral compounds obtained from the earth's crust. This distinction prevailed until Wohler, in 1828, prepared urea a characteristic organic compound from inorganic substances. Since then the barrier between the two classes of compounds has been completely removed. We now believe that compounds of carbon, whatever their source, are subject to the laws that govern all other compounds. The terms organic and inor- ganic are still used, though they have lost their original narrow meaning. Carbon forms a vast number of com- pounds which are related to each other, and which differ markedly from most compounds of other elements. It is convenient, therefore, to distinguish these compounds by the term organic and to study them under the comprehen- sive title of Organic Chemistry or the Chemistry of Carbon Compounds. Composition of Organic Compounds. The number of organic compounds is very large, but they contain only a few elements seldom more than four or five. Hydro- 405 406 Descriptive Chemistry. carbons, as already indicated, contain carbon and hydro- gen. Vegetable substances, typified by starch, sugar, and fruit acids, contain carbon, hydrogen, and oxygen. Ani- mal substances, like hah", albumen, gelatine, and muscle generally contain nitrogen as well as carbon, hydrogen, and oxygen ; some also contain sulphur or phosphorus. Artificial organic compounds, like dyestuffs, may contain any element, especially chlorine, iodine, and metals. The number and complexity of organic compounds is due to several facts already mentioned in a previous chapter, (i) Atoms of carbon have power to unite with themselves. (2) Atoms of different elements can be intro- duced into carbon compounds. Sometimes these atoms are simply added, sometimes they replace other atoms, thus producing an endless number of addition and substi- tution products. (3) The same number of atoms may arrange themselves differently, thereby producing isomeric compounds having different properties. To these princi- ples, which should be reviewed until firmly grasped, must be added another. (4) Organic compounds contain radi- cals. These radicals are analogous to hydroxyl (OH) and ammonium (NH 4 ), and like these radicals they exist only in combination. They act like single atoms and enter unchanged into a number of organic compounds. The radical C 2 H 5 is called ethyl. It is present in many organic compounds, and its presence in ordinary alcohol gives rise to the scientific name, ethyl alcohol. Methyl (CH 3 ) is another important radical, and phenyl (C 6 H 5 ) is especially common in the benzene series of organic com- pounds. Structure of Organic Compounds. An extensive study of the properties of organic compounds has revealed many facts about their constitution, i.e. the structure of their Some Common Organic Compounds. 407 molecules. Little or nothing, of course, is known about the shape, size, etc., of molecules, but much is known about the grouping of atoms and of radicals in the mole- cules. These facts, which are ascertained by experiment and are often too complex to be expressed briefly, may be represented by suitable formulas. The ordinary or empiri- cal formula of alcohol is C 2 H 6 O. But this formula tells nothing about the relation these atoms bear to each other, nor whether all the hydrogen atoms act alike. Experiment proves, however, that (i) one hydrogen atom acts differ- ently from the other five, and (2) one hydrogen atom is always associated with the oxygen atom in chemical changes. Hence, the formula C 2 H 5 . OH expresses more fully these facts. Such a formula is called a rational or constitutional formula. Sometimes constitution is ex- pressed by a graphic formula. Thus methane and ethane have the graphic formulas H H H I I I H C H H C C H I I I H H H Methane Ethane In these diagrams the single lines represent a valence of one nothing else, and the number of lines connected with each atom must be equal to the valence of the ele- ment in the compound. The lines are sometimes called bonds or links, but they are not intended to represent at- traction or any other force. Nor do they represent space relations. In the case of methane, they mean that the four hydrogen atoms bear the same relation to the single carbon atom. In the case of ethane, they mean the same, 408 Descriptive Chemistry. but they also indicate that the two carbon atoms are joined. The graphic formula of ethyl alcohol is H H I I H C C O H I I H H This is not an arbitrary arrangement ; the facts mentioned above necessitate this general arrangement. Additional illustrations of this subject will be given, as different compounds are discussed. Classification of Organic Compounds. Organic com- pounds are divided and subdivided into many classes for purposes of study. Only the most common organic compounds can be considered in this book. These are members of the following groups: (i) Hydrocarbons, (2) Alcohols, (3) Aldehydes, (4) Ethers, (5) Acids, (6) Ethe- real salts, (7) Fats, glycerine, and soap, (8) Carbohydrates, (9) Benzene and its derivatives. Some compounds are so closely related that they really belong to several of these groups, while a few cannot strictly be put in any of them. HYDROCARBONS. Three of these compounds of carbon arid hydrogen have been fully considered in Chapter XV. 7 The chief facts and fundamental principles recorded there may be profit- ably reviewed at this point. Other hydrocarbons will be discussed under Benzene (see below). ALCOHOLS. Alcohols are compounds of carbon, hydrogen, and oxy- gen. Ordinary or ethyl alcohol is the best known member Some Common Organic Compounds. 409 of this group. It is usually called simply alcohol. There are many alcohols analogous to ethyl alcohol, but the only other important one is methyl alcohol. The alcohols may be regarded as hydroxides of certain radicals, e.g. ethyl, methyl, propyl, etc. 1 For example, ethyl alcohol is ethyl hydrox- ide, and may be considered as formed by replacing one hydrogen atom of ethane (C 2 H (i ) by one hydroxyl group (OH). Again, alcohols are analogous to metallic hydroxides, in which the metal is replaced by a radicals- Ethyl Hydroxide Sodium Hydroxide Alcohols and metallic hydroxides have some properties in common. Thus, both form salts with acids. With acetic acid, sodium hydroxide forms sodium acetate, while alcohol forms ethyl acetate (see Ethereal Salts). Methyl Alcohol, CH 3 .OH, is a colorless or slightly yellowish liquid, much like ordinary alcohol. It boils at about 66 C, and burns with a pale flame which de- posits no soot. It intoxicates, and if concentrated is poisonous. It mixes with water in all proportions. It is cheaper than ethyl alcohol, and is used as a solvent for fats, oils, and shellac, and in the manufacture of varnishes and dyestuffs. Methyl alcohol is often called wood alco- hol or wood spirit, because it is one of the liquid products obtained by the dry distillation of wood (see Charcoal). Ethyl Alcohol, C 2 H 5 . OH, is a colorless, volatile liquid, having a burning taste and a pleasant odor. It is lighter than water, its specific gravity being about 0.8. It boils at 78.3 C., and does not freeze until at 130.5 C. Be- cause of its very low freezing point, it is used in ther- 1 The names of these and similar radicals are derived from the correspond- ing hydrocarbon. Thus, the word methyl comes from methane, ethyl from ethane, propyl from propane. 4i o Descriptive Chemistry. mometers designed to record temperatures below 40 C. (the freezing point of mercury), as in Arctic explorations. Its harmful effect on the human system need not be dis- cussed. Alcohol mixes with water in all proportions. The ordinary commercial variety contains from 50 to 95 per cent of alcohol. Pure or absolute alcohol is obtained by removing the remaining water with lime. Proof spirit contains about 50 per cent of alcohol. Methylated spirit contains 90 per cent ethyl and 10 per cent methyl alcohol; it is often used as a cheap substitute for ordinary alcohol, but it cannot be used as a beverage on account of the dis- agreeable taste imparted by the methyl alcohol. Alcohol is an excellent solvent for gums, oils, and resins, and is therefore extensively used in the manufacture of varnishes, essences, extracts, tinctures, perfumes, and medicines. It is also used as an antiseptic, and as a source of heat in alcohol lamps. Many organic compounds, as ether and chloroform, are prepared from alcohol. Some vinegar is made from alcohol. In museums alcohol is used to pre- serve specimens. Alcohol may be prepared from ethane (see below), but it is manufactured by the fermentation of sugars. Fermentation is a general term for the chemical changes caused by ferments. The latter are usually minute living bodies, though some inorganic chemical sub- stances cause fermentation. The process and essential products vary with the nature of the ferment. The important kinds of fermenta- tion are alcoholic, acetic, and lactic, and the respective products are alcohol, acetic acid, FlG -77- and lactic acid. Alcoholic fermentation is Yeast cells. caused by ordinary yeast. Under the micro- scope, yeast has the form of slimy yellow chains of small, round cells (Fig. 77). When yeast is added to a solution Some Common Organic Compounds. 411 of glucose, or any other fermentable sugar, the yeast plants multiply rapidly. Air must be admitted, and the temperature should be 2O-3O C. The changes are numerous and complex, but the main products are alcohol and carbon dioxide, thus C 6 H 12 6 2C 2 H 6 + 2C0 2 Glucose Alcohol Carbon Dioxide The fermentation ceases as soon as the liquid contains about 14 per cent of alcohol. The solution is filtered and concentrated by distillation, until the distillate contains the desired per cent of alcohol. Commercial alcohol is made also from potatoes, grains, rice, beet root, molasses, and many other substances rich in sugar and starch. Ordinary or cane sugar must be boiled with acid before it will ferment. Wines, beers, and all alcoholic liquors are prepared by fermentation. Yeast is seldom added, however, because the ferment which brings about the change is in the air, upon fruits and vines. Wines are made from the juice of grapes ; beer is made from hops and malt (barley which has sprouted). Whisky, gin, brandy, rum, and cordials are called distilled liquors, and are manufactured by dis- tilling the liquid obtained by fermenting grains, molasses, fruit juices, and other substances containing sugar and starches. Hence, wine, beer, and similar liquors are essen- tially mixtures of alcohol and water. They differ mainly in their proportion of alcohol. The particular flavor is due to small quantities of different substances which are inten- tionally added, obtained from the raw materials, or formed by special processes of manufacture. Coloring matter is usually added, but sometimes it is extracted from the casks in which the liquor is stored. Beer contains from 3 to 7 412 Descriptive Chemistry. per cent of alcohol, wines from 6 to 20, rum, brandy, and whisky from 40 to 60 or more per cent. ALDEHYDES. Aldehydes are compounds of carbon, hydrogen, and oxy- gen. They are formed by the oxidation of alcohols. The two important members of this group are acetic aldehyde (or acetaldehyde) and formic aldehyde (or formaldehyde). Acetic Aldehyde, CH 3 . CHO, is usually called simply aldehyde. It is a colorless, very volatile liquid, and has a peculiar, suffocating odor. It is a vigorous reducing agent, and is sometimes used to precipitate silver, as a thin coating, from silver solutions. It is converted by oxi- dizing agents into acetic acid (hence its name, acetic aldehyde}. Alde- hyde is prepared by oxidizing alcohol with a solution of potassium (or sodium) dichromate and sulphuric acid. When a mixture of these three substances is gently warmed, the characteristic odor of aldehyde may be detected. The oxidation of alcohol consists simply in the removal of hydrogen, thus C 2 H,.OH + O = CH 3 .CHO + H 2 O Alcohol Aldehyde The word aldehyde emphasizes this fact, being a contraction of 0/cohol When chlorine is used to oxidize alcohol, part of the hydrogen is replaced by chlorine, and the compound CC1 3 . CHO is formed. This substance, called chloral, forms a hydrate (CC1 3 . CHO . H 2 O), which is used to induce sleep and relieve pain. When chloral is treated with an alkali, it is decomposed and chloroform (CHC1 3 ) is produced. The latter is a sweet liquid, and is used to produce insensibility in surgical operations. Chloroform is usually made by treating alcohol with bleach- ing powder. lodoform (CHI 3 ), which is analogous to chloroform, is a yellow solid, with a disagreeable smell, and is extensively used as a dressing for wounds. It protects the wound from the harmful action of germs. Formaldehyde, H . CHO, is a gas, but is used only in solution. It has a penetrating odor. The commercial solu- Some Common Organic Compounds. 413 tion sold as formalin contains 40 per cent of formaldehyde. It corresponds to methane and methyl alcohol, thus H H H I I I H-C-H H-C-O-H C = O I I I I H H H Methane Methyl Alcohol Formaldehyde With oxygen it forms formic acid (hence its name, see below). Large quantities of formaldehyde are used in the manufacture of dyestuffs and fuming nitric acid, as a food preservative, and a disinfectant. When used for the last purpose, the solution is vaporized in a special kind of lamp, and the vapors are conducted by a small tube into the room to be disinfected. It is one of the most convenient and efficient of all disinfectants, and is very generally used. I ETHERS. Ethers are compounds of carbon, hydrogen, and oxygen. They are analogous to the metallic oxides. They are formed by heating alcohols with sulphuric acid. Ordinary or ethyl ether is the best known member of this group. Ethyl Ether, C 4 H 10 O, is a colorless, volatile liquid, with a peculiar, pleasing taste and odor. It is lighter than water, its specific gravity being about 0.74. It boils at 35 C, and the vapor is very inflammable. The liquid should never be brought near a flame. It is somewhat soluble in water, and it also dissolves water to a slight extent. It mixes with alcohol in all proportions. It is a good solvent for waxes, fats, oils, and other organic com- pounds. Its chief use is as an anaesthetic, i.e. to render one insensible to pain in surgical operations. 414 Descriptive Chemistry. Ether is manufactured by distilling a mixture of ethyl alcohol and sulphuric acid in the proper proportions. Hence, the names, ethyl or sulphuric ether. Ethylsulphuric acid is first produced, thus C 2 H 5 .OH + H 2 S0 4 HC 2 H 5 SO 4 + H 2 O Alcohol Sulphuric Acid Ethylsulphuric Acid When more alcohol and the ethylsulphuric acid are heated together, ether is formed, and sulphuric acid is reproduced, thus, HC 2 H 5 SO 4 + C 2 H 5 . OH = (C 2 H 5 ) 2 O + H 2 SO 4 Ether The process is thus continuous, a small quantity of sulphuric acid serv- ing to transform a large quantity of alcohol into ether. Ethyl ether is ethyl oxide, (C 2 H 5 ) 2 O or C 2 H 5 . O . C 2 H 5 . ACIDS. Organic Acids are compounds of carbon, hydrogen, and oxygen. It is a large class of compounds divided into several series, one of the most important of which is the acetic or fatty series. Its best known member is acetic acid ; several of the higher members occur in fats and oils. These acids are closely related to hydrocarbons, alcohols, and alde- hydes, as may be seen by the following formulas : H H I I H-C-H H-C-(OH) I I H-C-H H-C-H I I H H Ethane Ethyl Alcohol Acetic Aldehyde Acetic Acid It is thus possible to pass from a hydrocarbon through a correspond- ing alcohol and aldehyde to an acid. The characteristic group of atoms in organic acids is COOH (or O = C - O - H), and is called carboxyl. H 1 C = O O=C-(OH) 1 1 H-C-H H-C-H 1 1 H H Some Common Organic Compounds. ,415 Acetic Acid, C 2 H 4 O 2 or CH 3 . COOH. This is the most common organic acid. It is manufactured on a large scale by the dry distillation of wood. The dark red watery distillate, which is called pyroligneous acid, con- tains about 10 per cent of acetic acid besides a small per cent of methyl alcohol and many other organic compounds. This distillate is neutralized with lime or sodium carbonate, and the acetate formed is then decomposed and distilled with hydrochloric or sulphuric acid. The acetic acid which condenses in the receiver may be further purified by dis- tilling it with potassium dichromate and then filtering through charcoal. Sometimes the pyroligneous acid is distilled without neutralizing ; the distillate is then dilute, impure acetic acid, known as wood vinegar. If sodium acetate, prepared as described above, is fused and then distilled with concentrated sulphuric acid, the product is a very concentrated acetic acid. It is called glacial acetic acid, because at about 1 7 C. it becomes an icelike solid. Commercial acetic acid is a water solution containing about 30 per cent of pure acetic acid. It is a colorless liquid, having a pleasant odor and a sharp taste. It is slightly heavier than water. It mixes with water and alco- hol in all proportions, and like alcohol is an excellent solvent for many organic substances. Recently, it has begun to replace alcohol as a solvent for many drugs. Acetic acid is used to prepare acetates, dyestuffs, and other organic compounds, medicines, white lead, and in the manufacture of vinegar. Vinegar is dilute, impure acetic acid. It is prepared by oxidizing dilute alcohol, the essential change being repre- sented thus C 2 H 6 4- 2 = C 2 H 4 2 + H 2 Alcohol Oxygen Acetic Acid Water 416 Descriptive Chemistry. The transformation is accomplished by fermentation. Two processes are used, (i) When beer, weak wines, or cider are exposed to the air, they slowly become sour, owing to the conversion of alcohol into acetic acid. The change is caused by the presence and activity of a ferment, known as mycoderma aceti, or " mother of vinegar." Strong wines and pure dilute alcohol do not become sour, because the ferment cannot live in such liquids. (2) In the "quick vinegar process," impure dilute alcohol is oxidized by ex- posing it to an excess of air. The operation is conducted in tall vats or casks filled with beechwood shavings soaked in strong vinegar (Fig. 78). Holes at the bottom and top allow air to enter and escape freely. The alcoholic solu- tion is introduced at the top, trickles through the shavings, and collects at the bottom. In its passage it comes in contact with the ferment and oxygen, and is partially con- verted into vinegar. The operation is repeated until Thus prepared, the vinegar lacks the flavor, odor, and color of cider vinegar, but these deficiencies are often artificially supplied. Vinegar is used chiefly as a condiment for the table and in making pickles and similar relishes. The constitution of acetic acid has been shown to correspond to the formula CH 3 . COOH. Its metallic salts are formed by substituting a metallic atom (or group) for the hydrogen of the group COOH. radical CH 3 remains unchanged. (See page 170.) FIG. 78. Apparatus for the prep aration of vinegar from impure, dilute alcohol. the change is complete. The Some Common Organic Compounds. 417 Acetates. Acetic acid is a monobasic acid, and forms a series of salts the acetates. They are prepared like other salts by the interaction of the acid and carbonates, hydroxides, metals, etc. The metallic acetates are usually crystallized solids, which readily yield acetic acid when treated with sulphuric or a similar acid. Most of them contain water of crystallization, and most are poisonous. Several acetates have useful applications. Sodium acetate, NaC 2 H 3 O 2 . 3 H 2 O, is a white crystallized solid, used in preparing pure acetic acid, and in the manufacture of dyestuffs. Lead acetate, Pb(C 2 H s O 2 ) 2 , is a white crystallized solid, used, in dyeing and in mak- ing a yellow pigment. Its sweet taste led to the common name of "sugar of lead. 11 Aluminium acetate, A1(C 2 H 3 O 2 ) 3 , is not known in the pure state, but an impure solution, known as " red liquor," is exten- sively used in dyeing and calico printing. Iron acetates are sold in solution as a complex black liquid, known as "iron liquor," which is used in dyeing black silks and cottons, and in calico printing (see Mordants). A complex copper acetate, 2 Cu(C 2 H 3 O 2 ) 2 + CuO, called verdigris, is used in making blue paint. Another complex acetate of copper and arsenic is Paris green ; it is used to kill potato bugs and other insects which injure vegetation. A few other acids in this series are interesting. Butyric acid C 4 H 8 O 9 , is the acid which gives the disagreeable odor to rancid butter. Stearic acid. C 18 H 3(J O 2 , and Palmitic acid, C 16 Ho 2 O 2 , are found as compounds in beef suet, mutton fat, butter, and other fats. Palmitic acid is also one of the essential compounds found in palm oil. These two acids are white solids, and are used to make stearin candles (see Fats, below). Other Organic Acids which are important are oxalic, lactic, malic, tartaric, and citric. Oxalic Acid occurs as a salt in rhubarb and sorrel. It is manufactured on a large scale by heating sawdust with potassium hydroxide, and treating the residue first with lime and then with sulphuric acid. Oxalic acid is a white solid, very soluble in water, from which it crystallizes with 4i 8 Descriptive Chemistry. two molecules of water of crystallization (C 2 H 2 O 4 . 2. H 2 O). It is very poisonous. It is dibasic and forms several use- ful salts. The acid and some of its salts decompose iron rust and inks containing iron, and are often used to remove such stains from cloth. The acid and its salts are also used in dyeing, calico printing, photography, in making dyestuffs, and as an ingredient of mixtures for cleaning brass and copper. Lactic Acid, C 3 H 6 O 3 , occurs in sour milk, being one product of the fermentation of the milk sugar. It is a thick, sour liquid, and is easily decomposed by heat. When sour milk is used in cooking, the " baking soda " and lactic acid interact, producing soluble sodium lactate and carbon dioxide gas. Lactic acid and its salts are used as medicines, in beverages, and as a substitute for more expensive acids in dyeing and calico printing. Malic acid, C 4 H 6 O 5 , is found free and as salts in apples, pears, cur- rants, gooseberries, rhubarb, grapes, and berries of the mountain ash tree. It is a white, crystalline solid. Tartaric Acid, C 4 H 6 O 6 , occurs as the potassium salt in grapes and other fruits. During the fermentation of grape juice, impure acid potassium tartrate is deposited in the casks. From this argol or crude tartar the acid itself is prepared by treating the raw product successively with chalk and sulphuric acid. Tartaric acid is a white crystal- lized solid, soluble in water and alcohol. It is used in dye- ing, and as one ingredient of Seidlitz powders. In these and similar powders it serves to decompose the other in- gredient which is a carbonate (see Sodium Bicarbonate). Tartaric acid is dibasic and forms two classes of salts. Purified acid potassium tartrate obtained from argol is commonly known as cream of tartar. It is extensively used in the manufacture of baking powders. These, as a rule, are essentially mixtures of cream of tartar Some Common Organic Compounds. 419 and sodium bicarbonate, HNaCO 3 . When moistened by dough, the baking powder dissolves, the two ingredients interact and liberate car- bon dioxide as the main product. This gas bubbles slowly through the dough, thereby puffing it up and making it porous (see Sodium Bicarbonate). Tartar emetic is a tartrate of potassium and antimony. It is used as a medicine and to some extent in dyeing. Citric Acid, C(;H 8 O 7 , occurs abundantly in lemons and oranges, and in small quantities in currants, gooseberries, and raspberries. It is a white, crystallized solid, very soluble in water. The taste is sour, but pleasant. The acid and its magnesium salt are used as medicines. The acid itself is used in calico printing. Citric acid is tribasic. ETHEREAL SALTS. Ethereal Salts or Esters are compounds of carbon, hy- drogen, and oxygen closely related to alcohols and organic acids. Thus, when ethyl alcohol, acetic acid, and concen- trated sulphuric acid are mixed and warmed, ethyl acetate is formed. The essential change is represented thus ^ C 2 H 5 .OH +CH 3 .COOH = CH 3 .COOC 2 H 5 + H 2 O Ethyl Alcohol Acetic Acid Ethyl Acetate Water The sulphuric acid serves to absorb the water. Ethyl acetate has a pleasant, fruitlike odor, and its formation in this way is a simple test for alcohol or acetic acid. Ethyl acetate is analogous to sodium acetate, i.e. the organic salt contains the radical ethyl while the metallic salt con- tains sodium. The fatty acids, as well as those of other series, form many ethereal salts of special interest. Some occur naturally in fruits and flowers, and in many cases give the flavor and fragrance. Others are prepared artifi- cially and used as the basis of cheap flavoring extracts, perfumery, and beverages. Ethyl butyrate has the taste and fragrance of pineapples, amyl acetate of bananas, amyl valerate of apples. 420 Descriptive Chemistry. FATS, GLYCERINE, AND SOAP. General Relations. Natural fats and oils are essentially mixtures of stearin, palmitin, and olein. Beef and mutton fat are chiefly stearin, lard is mainly palmitin and olein ; while oils, such as olive oil, are largely olein. Stearin and pal- mitin are solids at the ordinary temperature, but olein is a liquid. These three compounds stearin, palmitin, and olein are ethereal salts of their corresponding acids and the alcohol, glycerine. They are analogous to ethyl acetate. The radical of glycerine is glyceryl, C 3 H 5 . Thus, stearin is glyceryl stearate, palmitin is glyceryl palmitate, and olein is glyceryl oleate. Natural fats and oils, therefore, are mixtures of these and similar ethereal salts. Fats are sometimes called glycerides. Glycerine is a triacid alcohol containing three hydroxyl (OH) groups. Like ordinary alcohol, it interacts with the fatty acids and forms ethereal salts. The latter, as we have just learned, are the fats. Now when fats are heated with very hot steam or with sul- phuric acid, the fats themselves are changed into glycerine and the corresponding acids. Thus, with stearin, the change is (C 17 H 35 . C0 2 ) 3 C 3 H 5 + 3 H 2 = C 3 H 5 (OH), + 3 C ir H,, . COOH Stearm Glycerine Stearic Acid But if fats are boiled with sodium hydroxide or a simi- lar alkali, glycerine and an alkaline salt of the correspond- ing acid are formed. Soap is a mixture of such alkaline salts. In a few words, the general relations are these: (i) fats are ethereal salts. (2) Treated with steam or acid, fats form glycerine and fatty acids. (3) Treated with alka- lies, fats form glycerine and soap. Natural Fats and Oils are often complicated mixtures. The solid fats, as already stated, are rich in stearin and Some Common Organic Compounds. 421 palmitin. Tallow is chiefly stearin, but human fat and palm oil are largely palmitin. The soft and liquid fats and oils contain considerable olein, as a rule. The proportion of olein determines the consistency of the fats and oils. Thus, Olive oil contains about 72 per cent of olein (and a similar fat) and 28 per cent of stearin and palmitin. The specific character of many fats and oils is due mainly to the presence of a small proportion of certain fats. These fats correspond to uncommon acids in the fatty, oleic, and other series. Butter, for example, consists mainly of the fats corresponding to the following acids : palmitic, stearic, oleic, butyric, capric, and caproic. The last three with traces of other substances give butter its pleasant flavor. Oleomargarine and other substitutes for butter resemble real butter very closely in composition. Artificial butter, however, lacks the flavor of the real butter, but it is " prob- ably just as nutritious, although perhaps not quite so easily digested." The lack of flavor noticed in artificial butter is due to the absence of the fats corresponding to the acids of low molecular weight. Cottolene is a mixture of beef fat and cotton-seed oil ; it is used as a substitute for lard. Glycerine (C 3 H 8 O 3 or C 3 H 5 .(OH) 3 ) is a thick, sweet liquid. It mixes readily with water and with alcohol in all proportions, and absorbs moisture from the air. Heated in the air, it decomposes and gives off irritating gases, like those produced by burning fat. Glycerine is used to make nitroglycerine (see below), toilet soaps, printers' ink rolls ; it is also used as a solvent, a lubricator, a preservative for tobacco and certain foods, a sweetening substance in certain liquors, preserves, and candy ; as a cosmetic ; and, owing to its non-volatile and non-drying properties, it is used as an ingredient of inks and oils. 422 Descriptive Chemistry. Glycerine is a by-product in the manufacture of soap, or it is made directly by decomposing fats with steam under pressure or with lime. Ail these methods involve the chemical change described above, viz. the decomposition of an ethereal salt (the fat) into the corresponding alcohol (glycerine) and a mixture of fatty acids. By skillful treatment the glycerine is freed from water and impurities. The mixture of fatty acids is made into the so-called "stearin" candles. As already stated, glycerine is an alcohol, and for this reason it is often called glycerol. When treated with a mixture of concentrated nitric and sulphuric acids, it forms an ethereal salt commonly known as nitroglycerine (C 3 H 3 (ONO 2 ) 3 ). This is a yellow, heavy, oily liquid. It is the well-known explosive, and is also an ingredient of some other explosives. When kindled by a flame, it burns without explosion ; but if struck by a hammer or heated suddenly by a percussion cap, it ex- plodes violently. Nitroglycerine is used in blasting ; but since it is dan- gerous to handle and transport, it is usually mixed with some porous substance, such as infusorial earth, fine sand, or even sawdust. In this form it is called dynamite. Soap, as already stated, is a mixture of alkaline salts of organic acids, mainly stearic and palmitic acids. Soap is made by boiling fats with sodium hydroxide or potassium hydroxide. This process is called saponification. Sodium hydroxide produces hard soap, consisting chiefly of sodium palmitate, sodium stearate, and sodium oleate. Potassium hydroxide produces soft soap, which is mainly the corre- sponding potassium salts. The chemical change, as already stated, consists in tr e transformation of an ethereal salt (fat) into glycerine and an alkaline salt. In the case of pure stearin (glyceryl stearate) the change may be repre- sented thus C 3 H 5 (C 17 H 35 . C0 2 ) 3 + 3NaOH - 3 C 17 H 35 . CO 2 Na + C 3 H 5 (OH) 3 Stearin Sodium Sodium Glycerine Hydroxide Stearate The fats used in soap making vary with the soap. Tal- low, lard, palm oil, and cocoanut oil make white soaps. Some Common Organic Compounds. 423 Bone grease or house grease, together with tallow, palm oil, cotton-seed oil, and rosin, make yellow soaps. Olive oil is used for making castile soap. In the^cold process the calculated amounts of alkali and fat are allowed to interact, first in a large tank and then in a box called a " frame." By this process the glycerine and excess of alkali are left in the soap. Most soaps are made by the boiling process. The fat and alkali are boiled in a huge kettle. This operation produces a thick, frothy mixture of soap, glycerine, and alkali. At the proper time, salt is added, thereby causing the soap to separate and rise to the top. The liquid beneath is drawn off, and from it glycerine is extracted. The soap is often boiled again with rosin or cocoanut oil ; then purified by washing, mixed, if desired, with perfume, coloring matter, or some filling material (such as sodium silicate, sand, borax), cooled in "frames," cut, and dried. Most soaps contain water. This really assists their cleansing action. The latter is believed to be due to the free alkali formed by the decomposi- tion of the soap when dissolved. CARBOHYDRATES. Carbohydrates are compounds of carbon, hydrogen, and oxygen. This is a large group, and the most important members are the sugars, starches, and cellulose. The term carbohydrate is applied to these compounds because they contain hydrogen and oxygen in the proportion to form water. They were once regarded as hydrates of carbon, or carbon hydrates a view which is incorrect and misleading. Sugars. The popular term sugar means almost any sweet substance found in fruits, nuts, vegetables, sap of trees, etc., though it is usually restricted to the ordinary white sugar obtained from sugar cane and sugar beet. Chemically, there are many sugars, each having a defi- nite constitution. The most important is ordinary sugar, which is also called cane sugar, sucrose, and saccharose. Another important sugar is glucose. 424 Descriptive Chemistry. Cane Sugar, C 12 H 22 O n , is widely distributed in nature, being found in the sugar cane, sugar beet, sugar maple, Indian corn, sorghum, most sweet fruits, many nuts, blos- soms of flowers, and honey. The main source of cane sugar is the sugar cane and sugar beet. Saccharose, or ordinary sugar, is a white, crystallized solid. Rock candy is highly crystallized sugar. It is solu- ble in water, but only sparingly soluble in alcohol. Heated to 160 C, sugar melts, and on cooling forms a pale yellow colored mass, called barley sugar. Heated to about 200 C., it is changed into water and a brown mass, called caramel, which is used to color liquors, soups, etc. If sugar is heated with sulphuric acid, it is changed into a black mass, which is mainly carbon ; several gases are also produced, such as steam, carbon dioxide, and sulphur dioxide. Cane sugar does not ferment. The manufacture of Cane Sugar from sugar cane and sugar beets involves two main operations: (i) the preparation of raw sugar and (2) its purification or refining, (i) In the preparation of raw sugar from sugar cane the juice is extracted from the cane by crushing the latter between heavy iron rollers. The liquid is then clarified as soon as possible by boiling it with a little lime, removing the scum which contains much of the impurity, and finally filtering the liquid through bags or a filter press. The purified juice is next evaporated until the cane sugar begins to crystallize from the cooled liquid. Formerly the evaporation was accomplished in an open pan, and is now in some localities, but usually a vacuum kettle is used. The crystals are next separated from the liquid by allowing the latter to drip out, or more commonly by whirling it out in a centrifugal machine. The solid product is called muscovado, raw or brown sugar. The thick liquid is the familiar molasses. There are several grades of each product. The preparation of raw sugar from sugar beets resembles the method used for sugar cane. The washed beets are reduced to a pulp, or cut into slices, and then treated with water. The sugar dissolves in the water. The solution is clarified, evaporated, and separated by pro- cesses much like those applied to cane-sugar solutions. The raw sugar Some Common Organic Compounds. 425 can scarcely be distinguished from cane sugar. The molasses is unfit for table use, though considerable sugar is extracted from it by means of strontium hydroxide (see Strontium Hydroxide). (2) Raw sugar is usually dark colored, and must be refined before it is suitable for most uses. The refining of sugar consists in (a) purification, and (<) recrys- tallization. () The raw sugar is purified by first dissolving it in huge tanks. Air is blown in to agitate the heated solution, blood and other substances are often added to entangle the impurities, and lime is also added to precipitate and gather the impurities into a scum or clot. The colored liquid is next filtered, first through cloth bags and then through animal charcoal, from which it drips as a perfectly clear liquid, (b} The filtered sirup is now evaporated in a large vacuum kettle. When a sample shows that the evaporation has reached the proper point, the liquid is run into tanks to crystallize. The crystals of sugar are separated from the sirup by centrifugal machines. The latter is boiled again or sold as sirup for the table. The crystals are dried in a heated tube called a granulator, so that each grain will be separate. Hence the name granulated sugar. The grains are sifted and packed in barrels for the market. Lactose, or sugar of milk, has the same formula as cane sugar, but its constitution and properties differ. It is obtained from milk. Its crystals are white, hard, gritty, less sweet than cane sugar ; they con- tain one molecule of water of crystallization. Sugar of milk is used in making homeopathic pills and certain kinds of foods for infants. Glucose is the name of a sugar and of a commercial mixture of glucose and several related substances. Glu- cose (dextrose or grape sugar, C 6 H 12 O 6 ) is found in many sweet fruits, especially in grapes. Old raisins are some- times coated with this sugar. It is often associated with levulose (fructose or fruit sugar) an is6meric compound (C 6 H 12 O 6 ). The two sugars are found, for example, in honey and in parts of some plants. Both sugars are formed from cane sugar by boiling it with a dilute acid. The chemical change may be represented thus C 12 H 22 O n + H 2 O = C 6 Hi 2 O 6 + C 6 H 12 O 6 Cane Sugar Glucose Fructose 426 Descriptive Chemistry. Both glucose and fructose ferment, forming alcohol and carbon dioxide (see Alcohol). The commercial mixture called "glucose" is prepared on a large scale by boiling starch with a dilute acid, usually sulphuric acid. The consistency and composition of the product vary with the details of manufacture. The liquid products are called " glucose " or "mixing sirup," while the solid product is known as " grape sugar " or " dex- trose." All contain more or less glucose and are about three fifths as sweet as sugar. But since they dissolve in water, and are cheaper than cane sugar, they are used extensively in the manufacture of candy, jelly, table sirups, etc. They are also added to wines and liquors, certain medicines, and many thick liquids in which their presence is harmless. In alkaline solutions, glucose is a strong reducing agent, and is used as such in dyeing with indigo. It also reduces an alkaline mixture of cop- per sulphate, known as Fehling's solution. When this solution is boiled with glucose, a reddish copper compound (cuprous oxide) is formed. The presence of sugar in solution is often shown in this way. Starch is widely distributed in the vegetable kingdom. It is found in wheat, corn, and all other grains, in pota- toes, beans, peas, and similar vegetables, and in large quantities in rice, sago, tapioca, and nuts. Many parts of plants contain starch, for example, the stalk, stem, leaves, root, seed, and fruit. The food value of vegetables de- pends largely upon the starch they contain. FIG. 79. Starch grains (magnified) wheat (left), rice (center), corn (right). Starch is a white powder, as usually seen. But under the microscope it is found to consist of a mass of oval Some Common Organic Compounds. 427 grains, varying somewhat with the source (Fig. 79). Starch is only very slightly soluble in water. But if heated with water, the grains swell and burst, partially dissolve, and form a solution which, when cold, becomes the familiar starch paste. Starch in solution is turned blue by iodine, and its presence in many vegetables and foods may be readily shown by grinding the substance in a mortar with warm water and adding a drop of iodine solution. Starch is prepared on a large scale chiefly from corn and potatoes. The operation is mainly mechanical, and consists in separating the starch from the fatty, nitrogenous, and mineral matters in the raw product. Immense quantities are consumed as food, in laundries, in finishing cloth and paper, in making glucose, and as a paste. The composition of starch, according to some authorities, corre- sponds to the formula C 6 H 10 O,, but its formula is still being investigated. Dextrin is a sticky solid formed from starch by heating it to 2OO-25O C. or by treating it with dilute acids. It is soluble in water and forms a sticky solution. Commer- cial dextrin or British gum is a mixture of dextrin and similar compounds. Mucilage contains dextrin. Large quantities are used as the gum for the backs of postage stamps, and for sticking the colors to the cloth in calico printing. Dextrin is sometimes regarded as an intermediate product between starch and dextrose. Its composition, according to some authorities, corresponds to the formula C 12 H 20 O 10 , but the statement made about the composition of starch also applies to dextrin. Bread. Wheat flour contains about 70 per cent of starch. The re- mainder is chiefly water and gluten in nearly equal proportions, though small quantities of mineral matter, dextrin, and other fermentable sub- stances are present. In making bread, flour, milk or water, and a little yeast are thoroughly mixed into dough, which is put in a warm place to rise, Fermentation begins at once. The yeast changes the ferment- 428 Descriptive Chemistry. able substances into alcohol and carbon dioxide. The gases, in trying to escape, puff up the dough, which literally rises and becomes light and porous. When the dough is baked, the heat kills the yeast, and fer- mentation stops ; but the alcohol, carbon dioxide, and some water escape and puff up the mass still more. The heat, however, soon hardens the starch, gluten, etc., into a firm but porous loaf. Cellulose (C 6 H 10 O 5 ) n is widely distributed in the vegetable kingdom. The framework of all vegetables is cellulose. It is thus analogous to the bones of animals. Wood, cotton, linen, and paper are largely cellulose. Pure cellulose is a white substance, insoluble in most liquids, but soluble in a mixture of ammonia and copper oxide. Concentrated sul- phuric acid dissolves it slowly ; and if the solution is di- luted and boiled, the cellulose is changed into a mixture of glucose and dextrin. By this operation, wood could be made into a sugar and then into alcohol ; but the method would be too expensive to use on a large scale. Sulphuric acid of a certain strength, if quickly and properly applied to paper, changes it into a tougher form called parchment paper. The latter is often substituted for animal parchment (e-g. sheepskin), and has a variety of uses. Cellulose has properties resembling those of alcohol. Thus it inter- acts with acids and forms ethereal salts. With nitric acid it forms cellu- lose nitrates, just as glycerine forms glycerine nitrates (see Nitroglyce- rine). The cellulose nitrates are the basis of smokeless gunpowders. One of the cellulose nitrates is gun cotton. It looks like ordinary cotton, and may be spun, woven, and pressed into cakes. It burns with a large flame if unconfined ; but when ignited by a percussion cap or when burned in a confined space, gun cotton explodes violently- It is used in blasting. Other cellulose nitrates are known. Their solution in a mix- ture of alcohol and ether is called collodion. When poured or brushed upon a glass plate or the skin, the solvent evaporates, leaving behind a thin film. It is used in preparing certain photographic material and as a coating for wounds. The " new skin " liquid recently offered for sale is mainly collodion. It protects wounds from dusty, impure air, and thereby facilitates the healing. A mixture of camphor and cellulose ni- Some Common Organic Compounds. 429 trates is called celluloid. It is easily molded into various shapes. The white celluloid is made into collar buttons, and the colored varieties are made into toilet articles and ornaments. Celluloid smells of camphor, can be lighted with a match, and burns freely with a smoky flame. Paper is chiefly cellulose. Formerly it was made from various kinds of rags ; but now it is made almost entirely from wood, especially the paper used for newspapers and cheap books. The best paper, such as writing paper, is still made from linen rags. In making paper from wood, the latter is reduced to a pulp, which is washed, spread on a frame or an endless wire gauze, dried, and pressed. The pulp is prepared by two processes, the mechanical and the chemical. Mechanical pulp is made by holding a stick of wood against revolving stone upon which water constantly falls. Chemical pulp is made by heating chipped wood with caustic soda, or with cal- cium acid sulphite (usually called bisulphite). The operation is con- ducted under pressure in huge tanks called digesters. Chemical pulp has longer and stronger fibers than mechanical pulp. The two kinds of pulp are often mixed. Most paper is loaded, that is, clay, gypsum, or other mineral matter is mixed with the pulp to give the paper body. Paper intended for printing or writing is sized, that is, the surface is coated with gelatine, rosin, or a similar substance to prevent the ink from spreading. Many kinds are also smoothed by passing them between heavy rollers. Blotting and tissue papers are not sized or loaded. BENZENE AND ITS DERIVATIVES. Introduction. The hydrocarbon benzene was mentioned in Chapter XV as the first member of an homologous series. In the same chapter coal tar was described as a black, complex liquid obtained as a by-product in the manufacture of illuminating gas. Now, coal tar is the chief source of benzene and some of its related compounds, while from benzene itself hundreds of derivatives have been prepared. Some are absolutely indispensable to man, but many have 430 Descriptive Chemistry. as yet merely scientific interest. Only the most important benzene compounds can be described in this book. Benzene, C 6 H 6 , is a colorless liquid, lighter than water, and has an odor suggesting coal gas. It burns with a luminous, smoky flame, owing to its richness in carbon. Ordinary illuminating gas owes its luminosity partly to benzene. It dissolves fats, resins, iodine, sulphur, and rubber. Benzene is sometimes called benzol. It should not be confused with benzine, which is a mixture of hydro- carbons derived from petroleum. Benzene is chiefly used in preparing its derivatives. The Constitution of Benzene has been carefully studied. For rea- sons too extended to state here, it is believed that in a molecule of benzene the carbon atoms are arranged in a ring. The structural for- mula is often written thus H I C / \ H-C C-H II I H-C C-H \ ^ C I H Benzene forms many derivatives. In all of them the six carbon atoms remain as a nucleus. No carbon atom can be removed from the benzene molecule without producing complete decomposition. But for the six hydrogen atoms, other atoms or radicals can be substituted. Hence, the almost infinite number of derivatives of benzene. Toluene, C H~ . CH 3 , is the second member of the benzene series. It may be regarded as methyl benzene ; or as phenyl methane, that is, methane (CH 4 ) in which one hydrogen atom is replaced by the radical phenyl (C 6 H 5 ). Toluene is obtained from coal tar, and resembles benzene in its properties. Nitrobenzene, C,.H 5 . NO 2 , is a yellow liquid formed by the inter- action of benzene and nitric acid. It is volatile, and has the odor of Some Common Organic Compounds. 431 bitter almonds. Although poisonous, it is used to produce the flavor of almonds in essences and perfumery. It is chiefly used, however, in the manufacture of aniline. Aniline, C 6 H 5 .NH 2 , is an oily liquid, slightly heavier than water. It is prepared on a large scale by reducing nitrobenzene with nascent hydrogen. From aniline are made many compounds known as aniline dyes. The starting point of these dyes is rosaniline, which is pre- pared by oxidizing a mixture of aniline and toluidine (C 6 H 4 . CH 3 . NH 2 ). Derivatives of rosaniline produce exceedingly brilliant colors in every variety of shade. Vast dyeing industries have risen since the value of coal tar was discovered (about 1860). Phenol, C 6 H 5 .OH, is a white crystalline solid. It has a smoky odor, is poisonous, and burns the skin. Coal tar is the source of phenol. A solution of phenol in water, popularly called carbolic acid, is used as a disinfectant. Derivatives of Phenol are important. Picric acid, or trinitrophenol (C (i H 2 (NO 2 )oOH), is a yellow crystalline solid used in dyeing silk yellow. Salts of picric acid the picrates are used in making explosives. Related to phenol are hydroquinone (C 6 H 4 (OH) 2 ) and pyrogallic acid (C (; Ho(OH) 3 ), which are used extensively as developers in photography. Acids, Aldehydes, and Ethereal Salts of the Benzene Series. The simplest acid is benzoic acid (C 6 H 5 . COOH). It occurs in certain balsams and gums. It is usually pre- pared from gum benzoin, and is a white crystalline solid with a fragrant odor. The corresponding aldehyde (ben- zoic aldehyde, C 6 H 5 .COH) is commonly called oil of bitter almonds. It is a fragrant liquid and is used to some extent as a flavoring substance. Salicylic acid (C 6 H 4 . OH. COOH) is a white crystalline solid, which is extensively used as a food preservative. Sodium salicylate 432 Descriptive Chemistry. is a common remedy for rheumatism. The corresponding aldehyde gives the fragrance to the wild flower known as meadowsweet; and methyl salycilate is the essential ingredient of the checkerberry. Naphthalene, C 10 H 8 , is a white, lustrous, crystalline solid obtained from coal tar. It has a penetrating, un- pleasant odor, and is used as a substitute for camphor under the name of " moth balls." Large quantities of naphthalene are used in making dyestuffs. Anthracene, C 14 H 10 , is a white crystallized solid, and, like naphthalene, is obtained from coal tar. It is one of the most important hydrocarbons, because from it alizarin is made. Alizarin is a valuable dyestuff, not only because it produces brilliant colors with different mordants, but also because most of these colors are fast, that is, they do not fade like many aniline colors. The Turkey red -so common on cotton goods, is produced by alizarin. Aliza- rin was formerly obtained from madder root, but now vast quantities are artificially prepared. Glucosides are substances occurring in many plants and vegetables. By the action of ferments they are changed into glucose and other substances that are benzene derivatives. Amygdalin, for example, is found in bitter almonds, cherry and peach kernels, and laurel leaves. The ferment emulsin, which also occurs in the plants, breaks up the amygdalin into oil of bitter almonds, hydrocyanic acid, and glucose. Tannin is also a glucoside. The tannins are a group of related com- pounds found in the leaves, bark, and other parts of the oak, hemlock, and pine trees, in sumach, gallnuts, tea, coffee, and numerous plants. Several acids have been obtained from tannins. The best known are gallic acid and tannic acid ; the latter is also often called simply tan- nin, and probably all tannins contain some tannic acid. Tannic acid changes into gallic acid according to the following equation C 14 H 10 9 + H 2 2C 7 H fi 5 Tannic Acid Gallic Acid Some Common Organic Compounds. 433 The formula of gallic acid may be written C (; H 2 (OH) 3 . COOH, thus showing its relation to benzene. Tannin, in whatever form, produces black compounds with iron salts. Its presence in tea, hemlock bark, etc., may be shown by the formation of a black precipitate upon the addition of ferrous sulphate. This property is utilized in making writing ink, though some kinds of ink are now made from aniline dyes. The tannin in oak and hemlock barks is used in tanning leather. When raw hides are soaked in solutions of tannin, the tannic acid changes certain substances in the skin into insoluble compounds, which remain in the hide, thereby converting it into the soft pliable form known as leather. Tannins are also used as mordants in dyeing silk, cotton, and linen. Alkaloids are complex compounds obtained from plants and vegeta- bles. The chief property is the power to produce marked physiological effects upon animals. All of them contain nitrogen, and resemble ammonia in having an alkaline reaction and in uniting directly with acids to form salts. Their commercial form is usually a salt. Many are used as medicines and drugs, although they are poisonous, especially if taken in large quantities. Theine or caffeine is the alkaloid obtained from tea and coffee. Nicotine comes from tobacco and is very poison- ous. Cocaine is obtained from the coca plant. One of its salts is used by surgeons and dentists to relieve pain. Quinine and cinchonine are extracted from the bark of the cinchona tree ; both are used as a remedy for fevers. Morphine is the chief alkaloid found in opium. The latter is the dried sap obtained from a certain part of the unripe poppy. Morphine in different forms is used to relieve pain and induce sleep. The two familiar medicines, laudanum and paregoric, contain prepara- tions of opium. Large doses of any form of opium may be fatal. EXERCISES. 1. How were organic and inorganic compounds' once defined ? Do they differ fundamentally ? What compounds are now included by the term organic? 2. What is the essential element in organic compounds ? What other elements are often present ? 3. Give four reasons for the vast number of organic compounds. 4. Define an organic radical. Name three. 5. Define constitution. Illustrate it by the empirical, rational, and graphic formulas of alcohol. 434 Descriptive Chemistry. 6. Name the nine important groups of organic compounds. 7. Review the general properties of hydrocarbons (see Chapter XV) . Name four hydrocarbons. 8. Define an alcohol. Discuss the constitution of alcohols. 9. Describe the preparation of methyl alcohol. State its properties and uses. Why is it called (a} methyl alcohol, and (b) wood alcohol ? 10. State (a) the properties, and (b) the uses of ethyl alcohol. n. What is (a) alcohol, (b) ethyl alcohol, (c) absolute alcohol, (d) methylated spirit, (e) proof spirit ? 12. What is fermentation ? What are ferments ? 13. Describe the preparation of alcohol. Discuss the preparation, composition, and properties of () wines and beers, and (<) distilled liquors. 14. What are aldehydes ? How are they related to alcohols and to hydrocarbons ? 15. Describe the preparation and properties of (a} acetic aldehyde, and (b) formic aldehyde. State the uses of the latter. What is its commercial name ? 1 6. What are ethers ? How are they related to alcohols ? 17. Describe the preparation, and state the properties and uses of ordinary ether. 1 8. What are organic acids ? Illustrate (by acetic acid) their rela- tion to hydrocarbons, alcohols, and aldehydes. 19. Describe the manufacture of acetic acid. State (#) its properties, and () its uses. 20. What is (a) pyroligneous acid, () glacial acetic acid, (<:) wood vinegar, (d) commercial acetic acid ? 21. Discuss the composition of acetic acid. 22. What is vinegar ? Describe its manufacture. State its proper- ties and uses. 23. What are acetates ? State their general properties. Describe four, and state their uses. 24. Name three other acids (besides acetic) in the fatty acid series. Why is this series so called ? 25. State the occurrence, properties, and uses of (a*) oxalic acid, () lactic acid, (c) tartaric acid, (d} citric acid. Where is malic acid found ? 26. What is (a} argol, (b) crude tartar, (c) cream of tartar, (d) tar- tar emetic ? Some Common Organic Compounds. 435 27. Review baking powder (see Sodium Bicarbonate). 28. What are ethereal salts ? How are they formed ? Where are they found ? Describe ethyl acetate. Name three other ethereal salts and state their properties. 29. What is the test for (a) alcohol, and () acetic acid ? 30. State clearly the general relations of fats to glycerine and soap. 31 . Name the chief ingredients of fats and oils. What is (a) tallow, (b} butter, (c} oleomargarine, (d) stearin ? 32. Describe the preparation of glycerine. State its properties and uses. 33. Discuss the constitution of glycerine. State the properties and uses of (a) nitroglycerine, and () dynamite. 34. What is soap ? Describe its general method of manufacture. What is the chemistry of its manufacture ? What fats and alkalies are used in making soap ? Describe (a) the cold process, and (b) the boil- ing process of soap making. 35. What are carbohydrates? Why is this term used? Name several carbohydrates. 36. What are sugars ? Name several. 37. Discuss the distribution of cane sugar. State its properties. What is (a) cane sugar, (b) sucrose, (c} saccharose, (d) barley sugar, (e) caramel ? For what is the last used ? 38. Describe the preparation of raw sugar from (a) sugar cane, and (b) sugar beets. 39. Describe the refining of sugar. 40. What is (a) granulated sugar, {b} brown sugar, (c} molasses ? 41. What is the sugar of milk ? What is its scientific name ? For what is it used ? 42. What is the formula of glucose ? W T hat other names has glu- cose ? Where is glucose found ? What sugar is closely related to glucose ? How is glucose formed from cane sugar ? State the equation for the reaction. 43. How is commercial glucose prepared ? What is (a) commercial grape sugar, and (b} " glucose " ? State the properties and uses of commercial glucose. 44. Describe the test for sugar. 45. Discuss the distribution of starch. Describe starch. State its properties. What is the test for starch ? 46. How is starch prepared ? State its uses. 436 Descriptive Chemistry. 47. What is the simplest formula of starch ? How does it differ from the formula of () cane sugar, and ($) glucose ? 48. What is dextrin ? How is it prepared ? For what is it used ? 49. Discuss the chemistry of bread making. 50. What is cellulose ? Describe pure cellulose. State its properties. 51. What is (a) parchment paper, () gun cotton, (c) collodion ? 52. What is the chief constituent of paper ? Describe the manufac- ture of paper. 53. State the source of benzene. State its properties. What is (a) benzol, and (b) benzine ? 54. To what class of organic compounds does benzene belong? Why is it such an important compound ? 55. What is the chemical relation of benzene to (a) toluene, () nitrobenzene, (c) aniline, (d} phenol, (e) benzole acid ? 56. Describe nitrobenzene. What is its chief use ? 57. Describe aniline. How is it prepared ? For what is it used ? 58. Describe phenol. What is its source and use ? What is its common name ? 59. State briefly the relation of phenol to (a) picric acid, (]) pi- crates, (c} hydroquinone, (//) pyrogallic acid. What is the use of each ? 60. Describe briefly benzoic acid and benzoic aldehyde. 61. Describe salicylic acid. State the use of this acid. 62. Describe naphthalene. What is its popular name ? State its uses. 63. Describe anthracene. State its use. What is alizarin ? 64. What are glucosides ? Discuss (a} the occurrence, () the prop- erties, and (c) the uses of tannin. What is (a) ink, and () leather ? 65. What are alkaloids ? Name six. What is their chief property ? PROBLEMS. 1. Alcohol is 0.8 as heavy as water. What is the weight of 1200 cc. of alcohol ? 2. If 10 gm. of pure alcohol are burned, what weight of each product is formed ? (Equation is C 2 H (; O + 30, = 2 CO 2 + 3 H 2 O.) 3. Calculate the percentage composition of (#) alcohol (C 2 H 6 O), () acetic acid (C 2 H 4 O 2 , (c} cane sugar (C ]2 H 2 ,O n ). 4. Calculate the simplest formulas of the substances having the com- position : (a} carbon = 40, hydrogen = 6.67, oxygen = 53.33 ; () carbon = 15.8, hydrogen = 5.26, nitrogen = 36.84, sulphur = 42.1 ; (c) carbon = 54.55, hydrogen = 9.09, oxygen = 36.36. APPENDIX. 1. The Metric System. The fundamental unit of this system of weights and measures is the meter. It is the unit of length, and is 39.37 inches long. The meter and the other units have multiples and submultiples, which are designated by prefixes attached to the particular unit. The multiple prefixes are deca-, hecto-, and kilo-, equivalent respectively to 10, 100, and 1000. The submultiple prefixes are deci-, centi-, and milli-, which correspond respectively to o.i, o.oi, and o.ooi. The unit of weight is the gram. It is derived from the kilogram, which is the weight of a cubic decimeter of water at 4 C. A kilogram weighs about 2.2 pounds. Small weights are expressed in terms of the gram. Thus, the weight of an object weighing 2 grams, 2 centi- grams, and 5 milligrams is 2.025 grams. The unit of volume is the liter. It is equal to the capacity of the vessel containing a kilogram of water. A liter equals about one quart. The relation between the units, multiples, and submultiples is shown in the TABLE OF THE METRIC SYSTEM. LENGTH. WEIGHT. VOLUME. NOTATION. Kilometer Kilogram Kiloliter 1000. Hectometer Hectogram Hectolitefr 100. Decameter Decagram Decaliter 10. METER GRAM LITER I. Decimeter Decigram Deciliter O.I Centimeter Centigram Centiliter 0.01 Millimeter Milligram Milliliter O.OOI From this table it is evident that 10 milligrams equal I centigram, 10 centigrams equal I decigram, 10 decigrams equal i gram, and so on. 4.37 438 Descriptive Chemistry. The relation of the metric system to weights and measures in com- mon use is shown by the TABLE OF METRIC EQUIVALENTS. meter = 39.37 inches kilometer = 0.62 mile centimeter = 0.39 inch liter = 0.908 quart liter = 1.056 quart '(liq.) gram = 15.432 grains kilogram = 2.2 pounds (avoir.) metric ton = 2204 pounds inch mile cubic inch quart (liq.) pound (avoir.) ounce (avoir.) ounce (troy) grain (apoth.) 2.54 centimeters 1.6 kilometers 16.39 cubic centimeters 0.9465 liter 0.4536 kilogram 28.35 grams 31.1 grams 0.0648 gram The passage from the English to the metric system may be accom- plished by utilizing the TABLE OF METRIC TRANSFORMATION. To CHANGE MULTIPLY BY Inches to centimeters Centimeters to inches 2-54 0-3937 16.387 Cubic centimeters to cubic inches 0.061 28.31; Grams to ounces (avoir.) 0^0353 0.0648 Grams to grains J 543 The customary abbreviations of the common denominations are meter, m. decimeter, dm. centimeter, cm. liter, 1. kilogram, kg. or Kg. decigram, dg. cubic centimeter, cc. milligram, mg. centigram, eg. The ^referable abbreviation for gram is gm. The same abbreviation is used for singular and plural, e.g. I m., 4 gm., 3 cm., 50 cc. A convenient relation (true only in the case of water) to remember is i I = i kg. = i cu. dm. = 1000 cc- ^ 1000 gm. = 2,3 lb. Appendix. 439 PROBLEMS. 1. What is the abbreviation of gram, centigram, liter, meter, cubic centimeter, centimeter, decimeter, milligram ? 2. Express (a) i liter in cubic centimeters, () 2 1. in cc., (c) i meter in centimeters, (d) 250 cm. in dm., (e) i kg. in grams, (/) 250 gm. in mg. 3. Add 2 kg., 5 dg., 2 eg., 4 gm., and 7 mg., and express the sum in grams. 4. How many cc. in a liter ? 5. What is the weight in grams of (a) i liter of water, (<) 250 cc., (c) 500 cc., (d) 721 cc. ? 6. Express in grams (a) 721 kg., () 62 mg., (c) 245 eg., (d} 84 dg. 7. Express (a) 40 meters in inches, () 25 kilograms in pounds, (c} 54 grams in ounces, (d) 72 grams in grains, (e) 75 liters in quarts (liq.). 2. The Thermometer in scientific use is the centigrade. The boil- ing point of water on this thermometer is zoo, and the freezing point is o (Fig. 80). The equal spaces between these points are called degrees. The abbreviation for centigrade is C., and for degrees is . Thus, the boiling point of water is 100 C. Degrees below zero are always designated as minus, e.g. 12 C., means 12 degrees below zero. The thermometer in popular use is the Fahrenheit. On this instrument the boiling point of water is 212 and the freezing point is 32 above zero (Fig. 80). To change Fahrenheit degrees into the equivalent centigrade degrees, subtract 32 and multiply the remainder by f , or briefly C = f(F- 3 2). To change centigrade degrees into the equivalent Fahrenheit temperature, multiply by f and add 32 to the product, or briefly 212 FIG. 80. Ther- mometers. The point 273 C. is called absolute zero. Absolute temperature is reckoned from this point. Degrees on the absolute scale are found by adding 273 to the readings on the centigrade thermometer. Thus, 273 absolute is o C., 274 absolute is + 1 C., etc. 440 Descriptive Chemistry. PROBLEMS. 1. Change into Fahrenheit readings the following centigrade read- ings : (a) 60.5, (J) 40, (0 92, (<0 - 5> (') o, (/)ioo, () 860, (A) -40. 2. Change into centigrade readings the following Fahrenheit read- ings : (a) 207, (b) 1 80, (0 o, (W) -30, (*) 212, (7) 100, () -40, (X) 270. 3. Express the following centigrade readings in absolute readings : () o, (*) 24, (*) -13,00 - 26 - 3. Crystallization. Most substances in passing from a liquid or a gas into a solid assume a definite shape. This change is called crys- tallization, and the substances are said to crystallize or to form crys- tals. Crystals are produced by (i) evaporating a solution, (2) cooling a melted solid, or (3) cooling a vapor. Thus, salt crystals are formed by evaporating a salt solution ; sulphur crystals, by melting and then cooling sulphur, and iodine crystals, by heating iodine in a test tube. These methods are called, respectively, evaporation, fusion, and sub- limation. As a rule each substance has an individual crystal form by which it can be distinguished. Although there are thousands of different crys- tals, all belong to one of six classes or systems. This classification is based upon two assumptions: (i) all crystals contain certain lines called axes, and (2) the surfaces or faces are grouped around the axes in definite positions. The axes connect angles, edges, or faces, which are similarly situated on opposite sides of the crystal. The bounding planes or faces are arranged symmetrically around the axes, which also determine (by their lengths and relative positions) the positions of the bounding planes. For example, the cube has three equal axes at right angles to one another, and terminating in the center of each of the six bounding surfaces. The following is a brief description of the six 'systems of crystal- lization : FIG. 81. Isometric crystals (cube, octahedron, dodecahedron). Appendix. 441 (i) Isometric. This has three equal axes intersecting at right angles. The simplest forms are the cube, octahedron, and dodecahedron (Fig. 81). Substances crystallizing in this system are diamond, com- mon salt, alum, fluor spar, iron pyrites, and garnet. FIG. 82. Tetragonal crystals. (2) Tetragonal. This has three axes at right angles ; but one axis is shorter or longer than the other two, which are equal. The common forms are the prism, pyramid, and their combinations (Fig. 82). Tin dioxide and zircon form tetragonal crystals. FIG. 83. Orthorhombic crystals. (3) Orthorhombic. This has three unequal axes intersecting at right angles. Common forms are the prism, pyramid, and their com- binations (Fig. 83). Potassium nitrate, barium sulphate, topaz, and native sulphur crystallize in this system (see Fig. 49). FIG. 84. Hexagonal crystals. (4) Hexagonal. This has four axes : three are equal and intersect at 60 in the same plane ; the fourth is longer or shorter than the others 442 Descriptive Chemistry. and is at right angles to their plane. It is a complex system. Common forms are the prism, pyramid, rhombohedron, scalenohedron, and their combinations (Fig. 84). In this system are found quartz, calcite, beryl, corundum, and ice (see Figs. 5, 52, 61). (5) Monoclinic. This has three unequal axes : two cut each other obliquely, and the third is at right angles to the plane of the other two. Common forms are combinations of prisms. It is a complex system, but includes many substances, e.g. sulphur deposited by fusion, sodium carbonate, borax, gypsum, and ferrous sulphate (Fig. 85). FIG. 85. Monoclinic crystal. FIG. 86. Triclinic crystals. (6) Triclinic. This has three unequal axes, all intersecting at oblique angles. Common forms are complex combinations. Copper sulphate, potassium dichromate, boric acid, and several minerals form triclinic crystals (Fig. 86). 4. History and Biography. The biographical data and table given here will serve as a basis for this interesting branch of chemistry. Additional facts can be obtained from the historical books mentioned below (under " Reference Books ") . Arrhenius, Svante, 1859 . Swedish physicist. Contributor to modern theory of solution. Avogadro, Amadeo, 1776-1856. Italian chemist and physicist. Pro- posed in 1811 his hypothesis equal number of molecules in equal volumes of all gases at same temperature and pressure. Balard, Antoine Jerome, 1802-1876. French chemist. Discovered bromine in 1826. Becher, Johann Joachim, 1635-1682. German physician. Dis- covered few facts, but collected and explained writings of others. Believed in alchemy, but made no search for gold. Laid foundations of phlogiston theory. Appendix. 443 Bergman, Torbern, 1735-1784. Swedish chemist. Improved methods of chemical analysis. Believed in phlogiston. Studied min- erals and organic acids. Contributed much to the industrial develop- ment of Sweden. Intimate friend of Scheele. Berthollet, Claude Louis, 1748-1822. French chemist. Studied composition of ammonia, properties and nature of chlorine, hydrogen sulphide, and hydrocyanic acid. Explained chemical changes by " affinity." His discussion with Proust led to law of definite proportions. Berzelius, Johann Jacob, 1779-1848. Swedish chemist. Deter- mined many atomic weights. Introduced use of symbols. Discovered selenium, prepared silicon and several rare elements. Investigated law of multiple proportions, proposed dualistic theory and an electrochem- ical theory, improved experimental methods. Industrious investigator, prolific writer. Bessemer, Sir Henry, 1813-1898. English metallurgist. Devised, in 1856, Bessemer process of making steel. Black, Joseph, 1728-1799. Scotch chemist and physicist. Dis- covered carbon dioxide. Showed relation of this gas to carbonates of alkalies and alkaline earths. Opposed phlogiston theory. Teacher and friend of James Watt and Rutherford. Boyle, Robert, 1626-1691. English philosopher. Announced law of effect of pressure on gases. Studied air and water. Opposed to alchemy. Views anticipated present conception of constitution of matter. Laid foundation of qualitative analysis. Bunsen von, R. W. E., 1811-1899. German chemist. Studied blast furnace and developed gas analysis. Invented the burner, pho- tometer, and battery bearing his name. With Kirchhoff (about 1860) devised the spectroscope, and by it developed spectrum analysis and discovered rubidium and caesium ; improved the calorimeter ; studied chemical action of light. Cannizzaro, Stanislao, 1826 . Italian chemist. Revived Avoga- clro's hypothesis in 1858, and thereby led to revision of atomic weights. Cavendish, Henry, 1731-1810. English chemist. Discovered hy- drogen, determined specific gravity of gases, showed (i) solubility of calcium carbonate in water containing carbon dioxide, (2) formation of water by burning of hydrogen. Determined composition of the atmos- phere and of nitric oxide. Accepted phlogiston theory. He was parsimonious, eccentric, shy ; trained mathematician and electrician ; " the richest of the wise, and the wisest of the rich." 444 Descriptive Chemistry. Charles, Jacques Alex Cesar, 1746-1822. French physicist. Pro- posed law bearing his name. Courtois, Bernard, 1777-1838. French chemist. Discovered iodine in 1811. Dalton, John, 1766-1844. English chemist, physicist, and mathe- matician. Devised atomic theory. Discovered law of multiple propor- tions. " Dalton was often inaccurate as to facts, deficient in the details of chemical manipulations, and did not hold high rank as an experi- menter; but he was good at drawing conclusions and at stating generalizations, his aim being the establishment of general, underlying laws." (Venable.) Davy, Sir Humphry, 1778-1829. English chemist. Studied gases, demonstrated properties of nitrous oxide, determined composition of hydrochloric acid, studied iodine and chlorine, named latter. Isolated potassium, sodium, barium, calcium, and strontium by electrolysis, and studied action of electricity on water and on many other substances. Devised miner's safety lamp. "He was one of the most brilliant chemists the world has ever seen and the greatest England has pro- duced." Dewar, James, 1842 . English chemist. Pioneer in the lique- faction of gases by modern methods. (See Hydrogen.) Dulong, Pierre Louis, 1785-1838. French chemist and physicist. With Petit announced law of specific heats in 1819. Dumas, Jean Baptiste Andre, 1800-1884. French chemist. Deter- mined many atomic weights, gravimetric composition of water, compo- sition of air. Investigated many organic compounds. Devised a method of determining vapor density. Excellent teacher, careful editor, and faithful public servant. Faraday, Michael, 1791-1869. English chemist and physicist. Liquefied chlorine and other gases. Showed quantitative relation be- tween electric current and chemical changes, and developed electro chem- istry. Was Davy's assistant and successor in the Royal Institution. Popular lecturer, keen investigator, and ardent lover of science. Gay-Lussac, Joseph Louis, 1778-1850. French chemist and physi- cist. Announced law of gas volumes in 1808. Worked on cyanogen, iodine, halogen acids, alkaline oxides, isolation of boron. Improved methods of analyzing organic compounds. Was pupil of Berthollet. " Was a trained chemist, capable of most accurate analytical work, and possessing scientific acumen in a very high degree." (Venable.) Appendix. 445 Glauber, Johann Rudolph, 1604-1668. German chemist. Believed in alchemy. Discovered sodium sulphate, which even now bears his name. Suggested improvements in industrial chemistry. Graham, Thomas, 1805-1869. British chemist. Studied diffusion of gases, acids of phosphorus, water of crystallization, and dialysis. Developed idea of basicity of acids. Hofmann von, August Wilhelm, 1818-1892. German chemist. Studied organic chemistry exhaustively. Coal-tar industry arose largely from his work. Devised unique lecture apparatus, e.g. that for the electrolysis of water. Brilliant teacher, prolific investigator. Kirchhoff , Gustav Robert, 1 824-1 887. German physicist. With Bun- sen, devised spectroscope and founded principles of spectrum analysis. Lavoisier, Antoine Laurent, 1743-1794. French chemist. Over- threw phlogiston theory, explained combustion, contributed many facts to a large number of chemical topics. Devised foundation of chemical nomenclature. Interpreted experiments of other chemists. Efficient public servant. Regarded by many as the founder of modern chem- istry. Accused of appropriating public money and of " putting water in the people's tobacco," he was condemned by the infamous Robes- pierre, and publicly guillotined. Liebig von, Justus, 1803-1873. German chemist. Laid founda- tions of agricultural and organic chemistry. Eminent teacher. Mendeleeff, Dmitri Ivanovitch, 1834 . Russian chemist. An- nounced periodic law in 1868. Meyer, Lothar 1830-1895. German chemist. Contributed to estab- lishment of periodic law. Moissan, Henri, 1852 . French chemist. Isolated fluorine, devised and perfected electric furnace, prepared artificial diamonds, rare metals, and refractory compounds. Ostwald, Wilhelm, 1853 . German chemist. Contributor to modern theory of solution. Eminent teacher and prolific writer. Petit, Alexis Therese, 1791-1820. French physicist. (See Dulong.) Priestley, Joseph, 1733-1804. English chemist and theologian. Student of electricity, light, and gases. Discovered oxygen. Devised pneumatic trough. His political and religious views were so freely expressed that he was obliged to leave England. Came to America in 1795. Died at Northumberland near Philadelphia, Pennsylvania. Proust, Louis Joseph, 1755-1826. French chemist. Defended law of definite proportions in a long controversy with Berthollet. 446 Descriptive Chemistry. " One of the good results of this controversy was to bring about a defi- nition of compounds and mixtures, and a clear distinction between them. In course of it, also, Proust discovered the hydroxides, a class of compounds until then confused with the oxides." (Venable.) Prout, William, 1785-1850. English physician. Advanced in 1815 the hypothesis that the atomic weights of all elements are whole numbers. Ramsay, William, 1852 . English chemist. Discovered argon, helium, neon, krypton, and xenon. Rutherford, Daniel, 1749-1819. Scotch botanist and physician. Discovered nitrogen in 1772. Pupil of Black. Scheele, Carl Wilhelm, 1742-1786. Swedish chemist. Discovered chlorine, ammonia, manganese, baryta, many acids (organic and inor- ganic), and oxygen (independently of Priestley). Isolated and studied borax, glycerine, Prussian blue, microcosmic salt. Improved the methods of preparing many substances. Was very poor. Friend and companion of Bergman. Achieved marvelous results with simple appliances. Believed in phlogiston. Stahl, George Ernst, 1660-1734. German physician and chemist. Revived and extended Becher's ideas of combustion. Introduced the name phlogiston. Strongly advocated this theory. Successful teacher and writer. Stas, Jean Servais, 1813-1891. Belgian chemist. Determined accurately many atomic weights. Pupil of Dumas. Overthrew Prout's hypothesis. Van Helmont, Jean, 1577-1644. Dutch chemist. Studied gases, and discovered carbon dioxide. Had imperfect but introductory views on physiological chemistry, indestructibility of matter, and elements. Believed in the alkahest or universal solvent. Van't Hoff, Jacobus Hendricus, 1852 . Dutch chemist. Con- tributor to chemistry of space relations of atoms and to modern theory of solution. Wohler, Friedrich, 1800-1882. German chemist. Isolated alu- minium and beryllium. Worked on boron, silicon, and many organic substances. Discovered isomerism. Overthrew barrier between or- ganic and inorganic chemistry. Was fellow-worker with Liebig, pupil of Berzelius, and influential teacher of many famous chemists. Appendix. 447 CHRONOLOGICAL TABLE OF FAMOUS CHEMISTS. Greeks Galen Aristotle Geber Avicenna Albertus Magnus Roger Bacon 8th Century 978-1036 1193-1280 1214-1294 Middle Ages Raymond Lulli Basil Valentine 1235-1315 1394 I4th to i6th Cen- Paracelsus Agricola Libavius Van Helmont turies. 1493-1541 1494-1555 1540-1616 1577-1644 Glauber Boyle Becher Hooke lyth and i8th 1604-1668 1626-1691 1635-1682 1635-1702 Centuries. Mayow Stahl Boerhaave Hales 1645-1679 1660-1734 1668-1738 1677-1761 ENGLISH. Black Cavendish Priestley 1728-1799 1731-1810 1733-1804 Dalton Davy Faraday 1766-1844 1778-1829 1791-1867 i8th and igth FRENCH. Lavoisier Berthollet Proust Centuries. I743-J794 1748-1822 1755-1826 Gay-Lussac 1778-1850 SWEDISH. Bergman Scheele Berzelius 1735-1784 1742-1786 1779-1848 ENGLISH. Graham FRENCH. Dumas BELGIAN. Stas 1805-1869 1800-1884 1813-1891 igth Century. GERMAN. Wohler Liebig Bunsen 1800-1882 1803-1873 1811-1899 Hofmann 1818-1892 5. Atomic Weights. The following table of atomic weights is from the Journal of the American Chemical Society, Vol. XXV, No. I (January, 1903). 448 Descriptive Chemistry. TABLE OF ATOMIC WEIGHTS. S 1 ATOMIC WEIGHT. ELEMENT. > o> O=i6. Hi. APPROXIMATE. 1 Aluminium .... Al 27.1 26.9 27 Antimony Sb 120.2 "9-3 120 Argon A 39-9 39-6 Arsenic .... As 75-o"~ 74-4 75 Barium .... Ba 137-4 136.4 137 Bismuth .... Bi 208.5 206.9 Boron B n. 10.9 11 Bromine .... Br 79.96 79-36 80 Cadmium .... Cd 112.4 iii.6 Caesium .... Cs 133. 132. Calcium .... Ca 40.1 39-8 40 Carbon .... C I2.OO 11.91 12 Cerium .... Ce I4O. 139. Chlorine .... Cl 3545 35-18 35.5 Chromium .... Cr 52.1 5 1 -? 52 Cobalt Co 59-o 58.56 Columbium .... Cb 94- 93-3 Copper .... Cu 63.6 63.1 63.5 Erbium .... Er 166. 164.8 Fluorine . . . ' F 19. 18.9 19 Gadolinium Gd 156. 155. Gallium . Ga 70. 69-5 Germanium Ge 72.5 71.9 Glucinum .... Gl 9.1 9-3 Gold . . . . 4 Au 197.2 IQC.7 197 Helium .... He m TM 4- *-yj'/ 4- Hydrogen . . . . H 1.008 I.OOO 1 Indium .... In 114. 113.1 Iodine' I 126.85 125.00 127 Iridium .... Ir 193.0 191.5 Fe cqn rqr 56 Krypton .... Kr 3J';? 81.8 D3O 81.2 Lanthanum .... La 138.9 137.9 Lead Pb 206.9 2O^.^< 207 Lithium .... Li 7-03 OOJ 6.98 Magnesium .... Mg 24.36 24.18 24 Manganese .... Mn 55-0 54-6 55 Mercury . . . . Hg 200.0 198.5 200 Molybdenum Mo 96.0 95-3 1 Use these values in solving problems. Appendix. TABLE OE ATOMIC WEIGHTS (Continued}. 449 i / LTOMIC WEIGHT. ELEMENT. s O=i6. H-i. APPROXIMATE. 1 Neodymium Nd 143.6 142-5 Neon Ne 20. 19.9 Nickel Ni 58.7 58.3 Nitrogen N 14.04 13.93 14 Osmium .... Os 191. 189.6 Oxygen .... 16.00 15.88 16 Palladium .... Pd 106.5 105.7 Phosphorus .... P 31.0 30.77 31 Platinum .... Pt 194.8 193-3 195 Potassium .... K 39-15 38.86 39 Praseodymium . . . . Pr 140.5 1394 Radium .... Rd 225. 223.3 Rhodium .... ' Rh 103.0 IO2.2 Rubidium . . . Rb 854 84.8 Ruthenium .... Ru 101.7 100.9 Samarium ... Sm 150. 148.9 Scandium .... Sc 44.1 43-8 Selenium .... Se 79.2 78.6 Silicon Si 28.4 28.2 28 Silver Ag 107.93 107.12 108 Sodium . Na 23-05 22.88 23 Strontium ... Sr 87.6 86.94 Sulphur .... S 32.06 31.83 32 Tantalum .... Ta 183. 181.6 Tellurium .... Te 127.6 126.6 Terbium .... Tb 160. 158.8 Thallium .... Tl 204.1 202.6 Thorium .... Th 232.5 230.8 Thulium .... Tm 171. 169.7 Tin Sn 119. 118.1 119 Titanium .... Ti 48.1 '47.7 Tungsten .... W 184. 182.6 Uranium .... U 238-5 236.7 Vanadium .... V 51.2 50.8 Xe 128. 127. Ytterbium .... Yb 173.0 171.7 Yttrium .... Yt 89.0 88.3 Zinc Zn 654- 64.9 65 Zirconium .... Zr 90.6 89.9 1 Use these values in solving problems. 450 Descriptive Chemistry. 6. Reference Books and Supplementary Reading. The list of books given below will serve as the basis of a chemical library. The starred (*) titles indicate books intended for the teacher, though many parts of these books are not beyond the grasp of pupils. The library should contain at least numbers i, 5, 8, 10, 18, 20, 22, 24. Additional titles can be found in (i) List of Books in Chemistry, L. E. Knott Apparatus Co., Boston, Mass. ; (2) Smith and Hall's Teaching of Chemistry and Physics, p. 218; (3) NEWELL'S EXPERIMENTAL CHEM- ISTRY, APP. C, II. i. Text-Book of Inorganic Chemistry, Newth. Longmans, Green, & Co., 682 pp., $1.75. *2. General Inorganic Chemistry, Freer. Allyn & Bacon, Boston, 559 PP- #3- *3. Text-Book of Inorganic Chemistry, Holleman. John Wiley & Sons, 458 pp., $2.50. 4. Physical Chemistry for Beginners, Van Deventer. John Wiley & Sons, 154 pp., $1.50. 5. Chemical Theory for Beginners, Dobbin and Walker. The Macmillan Co., 236 pp., $ .70. *6. Introduction to Physical Chemistry, Walker. The Macmillan Co., 332 pp., $3. 7. The Birth of Chemistry, Rodwell. The Macmillan Co., 135 pp., $i. 8. Short History of Chemistry, Venable. D. C. Heath & Co., 172 pp., $i. 9. Faraday as a Discoverer, Tyndall. D. Appleton & Co., 171 pp., $i. 10. Short History of Natural Science, Buckley. D. Appleton & CO., 467 pp., $2. 11. Heroes of Science Chemists, Muir. Thomas Nelson & Son, 350 pp., $1.50. *I2. Essays in Historical Chemistry, Thorpe. The Macmillan Co., 582 pp., $4. 13. Humphry Davy, Thorpe. The Macmillan Co., 240 pp., $1.25. 14. John Dalton, Roscoe. The Macmillan Co., 216 pp., $1.25. 15. Michael Faraday, Thompson. The Macmillan Co., 308 pp., $1.25. *i6. Alembic Club Reprints, University of Chicago Press, $.40 each, (i) Experiments on Magnesia Alba. (2) Foundations of the Atomic Appendix. 451 Theory. (3) Experiments on Air. (4) Foundations of the Molecular Theory. (6) Decomposition of the Fixed Alkalies. (7) (8) Discov- ery of Oxygen. (9) Elementary Nature of Chlorine. (13) Early His- tory of Chlorine. *ij. Organic Chemistry, Remsen. D. C. Heath & Co., 426 pp., $1.30. 18. Outlines of Industrial Chemistry, F. H. Thorp. The Macmil- lan Co., 528 pp., $3.50. *I9- Practical Electro-Chemistry, Blount. The Macmillan Co., 374 pp., $3.25. 20. Chemistry in Daily Life, Lassar-Cohn. J. B. Lippincott Co., 336pp., $1.75. 21. The Soil, King. The Macmillan Co., 400 pp. 22. Story of a Piece of Coal, Martin. D. Appleton & Co., 165 pp., $.40. 23. Chemical History of a Candle, Faraday. Harper & Bros., 223 pp., $1.00. 24. Minerals and How to Study Them, E. S. Dana. John Wiley & Sons, 380 pp., $1.25. *25. Teaching of Chemistry and Physics, Smith and Hall. Long- mans, Green & Co., 384 pp., $1.50. 26. Story of Nineteenth-Century Science, Williams. Harper & Bros., 475 PP-> $2.50. 27. Stories of Industry, Vol. I, Chase and Clow. Educational Pub- lishing Co., Boston, 172 pp., $.40. Scientific American, Munn & Co., New York. $3.00 yearly; single copies, 8 cents. School Science, Ravenswood, Chicago, Illinois. $2.00 yearly (9 issues) ; single copies, 25 cents. Popular Science Monthly, The Science Press, New York. $3.00 yearly ; single copies, 25 cents. PART II EXPERIMENTS CONTENTS. PART II. (Numbers in parentheses indicate experiments.) PAGE INTRODUCTION . . . . . . ..... . 459 Bunsen Burner ; Heating ; Cutting and Bending Glass Tubing ; Filtering ; Constructing and Arranging Apparatus ; Manipula- tion ; Smelling and Tasting. PHYSICAL AND CHEMICAL CHANGES . . " . ' . . . . 467 Physical Change (i, 2, 3); Chemical Change (4). OXYGEN . . ... . . . ... . . 468 Preparation (5) ; Properties (6) ; Preparation from Mercuric Oxide (7)- HYDROGEN . . . -. . . . . . . ;. . 471 Preparation (8); Properties (9); Burning Hydrogen (10). WATER . . . . . . . . . . . 474 General Distribution (n); Tests for Impurities (12); Distillation (13); Solubility of Gases (14); Solubility of Liquids (15); Solu- bility of Solids (16); Supersaturation (17); Water of Crystal- lization (18); Efflorescence (19); Deliquescence (20); Solution and Chemical Action (21) ; Electrolysis (22) ; Water and Chloripe (23); Water and Sodium (24). THE AIR , 481 Composition (25); Water Vapor (26) ; Carbon Dioxide (27). ACIDS, BASES, AND SALTS 483 Properties of Acids (28) ; Properties of Bases (29) ; A Property of Salts (30); Nature of Common Substances (31); Neutralization (32). HEAT, LIGHT, ELECTRICITY, AND CHEMICAL ACTION ..... 485 Heat and Chemical Action (33, 34) ; Light and Chemical Action (35) ; Electricity and Chemical Action (36, 37). 455 456 Descriptive Chemistry. PAGE CHLORINE 486 Preparation (38) ; Properties (39) ; Bleaching Powder (40) ; Prep- aration of Hydrochloric Acid (41); Properties of Hydrochloric Acid Gas (42) ; Properties of Hydrochloric Acid (43); Tests for Hydrochloric Acid and Chlorides (44). COMPOUNDS OF NITROGEN .' . . . 490 Preparation of Ammonia (45); Properties of Ammonia Gas (46); Properties of Ammonium Hydroxide (47) ; Neutralization of Am- monia (48) ; Preparation of Nitric Acid (49) ; Properties of Nitric Acid (50); Test for Nitric Acid and Nitrates (51-52); Interaction of Sodium Nitrate and Sulphuric Acid (53) ; Nitric Acid and Metals (54); Nitric Acid and Copper, and Nitrogen Peroxide (55); Ni- trous Oxide (56); Sodium Nitrite (57); Aqua Regia (58). CARBON .,...'. . . 498 Distribution (59) ; Decolorizing Action (60) ; Deodorizing Action (61 ) ; Preparation of Carbon Dioxide (62) ; Properties of Carbon Dioxide (63) ; Interaction of Calcium Carbonate and Hydro- chloric Acid (64); Carbon Dioxide and Combustion (65); Car- bonic Acid (66) ; Carbonates (67) ; Detection of Carbonates (68) ; Acid Calcium Carbonate (69); Carbon Monoxide (70); Ethylene (71); Acetylene (72); Illuminating Gas (73); Combustion of Illuminating Gas (74); Bunsen Burner (75); Bunsen Burner Flame (76); Candle Flame (77); Kindling Temperature (78); Reduction and Oxidation (79). FLUORINE, BROMINE, AND IODINE . 511 Hydrofluoric Acid (80); Bromine (81); Potassium Bromide (82); Iodine (83); Tests for Iodine (84, 85); Detection of Starch (86); Potassium Iodide (87). SULPHUR . . '> . . . . ... . . . 514 Properties (88) ; Amorphous Sulphur (89) ; Crystallized Sulphur (90) ; Combining Power (91); Sulphur and Matches (92); Preparation of Hydrogen Sulphide (93); Properties of Hydrogen Sulphide Gas (94) ; Sulphides (95) ; Preparation of Sulphur Dioxide (96) ; Properties of Sulphur Dioxide Gas (97) ; Properties of Sulphurous Acid (98); Sulphuric Acid and Organic Matter (99); Test for Sulphuric Acid and Sulphates (100). Contents. 457 PAGE SILICON AND BORON 520 Silicic Acid (101); Borax Beads (102); Boric Acid (103). PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH 522 Properties of Phosphorus (104); Test for Arsenic (105); Test for Antimony (106); Test for Bismuth (107). SODIUM AND POTASSIUM . 522 Properties of Sodium (108); Sodium Hydroxide (109); Exercises; Properties of Potassium (no); Potassium Hydroxide (in); Potas- sium Carbonate (112); Exercises. COPPER, SILVER, AND GOLD , . , , 525 Properties of Copper (113); Tests for Copper (114); Interaction of Copper with Metals (115); Exercises; Preparation of Silver (116); Properties of Silver (117); Test for Silver (118); Exercises ; Test for Gold (119). CALCIUM, STRONTIUM, AND BARIUM . .... . . . 528 Tests for Calcium ( 1 20) ; Plaster of Paris (121) ; Exercises; Test for Strontium (122); Red Fire (123); Tests for Barium (124); Green Fire (125); Exercises. MAGNESIUM, ZINC, CADMIUM, AND MERCURY . . . / . 530 Properties of * Magnesium (126); Tests for Magnesium (127); Exercises ; Properties of Zinc (128); Tests for Zinc (129); Inter- action of Zinc and Metals (130); Exercises; Test for Cadmium (131); Properties of Mercury (132); Tests for Mercury (133); Mercurous and Mercuric Compounds (134); Exercises. ALUMINIUM . . . 532 Properties (135); Action with Acids and Alkalies (136); Aluminium Hydroxide (137); Tests (138); Alum (139). TIN AND LEAD 534 Properties of Tin (140); Action of Tin with Acids (141); Tests for Tin (142); Deposition (143); Properties of Lead (144); Tests for Lead (145); Deposition of Lead (146); Oxides of Lead (147); Compounds of Lead (148). CHROMIUM AND MANGANESE 537 Tests for Chromium (149); Chromates (150); Reduction of Chro- mates (151); Chromic Hydroxide (152); Chrome Alum (153); Tests for Manganese (154); Potassium Permanganate (155); Exercises. 458 Descriptive Chemistry. PAGE IRON, NICKEL, AND COBALT . 540 Properties of Iron (156); Ferrous Compounds (157); Ferric Com- pounds (158); Reduction of Ferric Compounds (159); Oxidation of Ferrous Compounds ( 1 60); Compounds of Iron (161); Exercises; Test for Nickel (162); Test for Cobalt (163). ORGANIC COMPOUNDS . . 542 Composition (164); Alcohol (165); Properties of Alcohol (166); Aldehydes (167); Ether (168); Acetic Acid (169); Vinegar (170); Test for Acetic Acid and Acetates (171); Acetates (172); Organic Acids (173); Ethyl Acetate (174); Soap (175, 176); Glycerine (177); Test for Sugar (178); Exercises ; Benzene (179). LABORATORY EQUIPMENT . . ... . . . . 549 Apparatus ; Chemicals ; Solutions. INTRODUCTION. 1. The Bunsen burner is used as the source of heat in most chem- ical laboratories (Fig. 87). It is attached to the gas cock by a piece of rubber tubing. When the gas is turned on, the current of gas draws air through the holes at the bottom of the tube, and this mixture when lighted burns with an almost colorless, /. e. non-luminous, flame. It is a hot flame and deposits no soot. The burner is lighted by turning on the gas full and holding a lighted match in the gas about 5 centimeters (2 inches) above the top of the burner. If the flame is not colorless, or nearly so, turn the ring at the bottom of the burner until the flame is a faint blue. The colorless flame should be used in all experiments unless the directions state otherwise, and should be from 5 to 10 centimeters (2 to 4 inches) high. The hottest part of the flame is near the top. FIG. 87. Bunsen burner. 2. Heating. The following directions should be observed in heating with the Bunsen burner : (1) The burner should always be lighted before any piece of apparatus is held over it, or before it is placed beneath a wire gauze which supports a dish or flask. (2) Glass and porcelain apparatus should not be heated when empty nor over a bare or free flame even if they contain something unless directions so state. Vessels requiring a support should be placed on a wire gauze which stands on the ring of an iron stand, and heated grad- ually from beneath. Hot vessels should be heated and cooled gradu- ally ; if removed from the gauze while hot, they should be placed on a block of wood or piece of asbestos board never on a cold surface. (3) Many experiments require the heating of test tubes. These tubes should be dry on the outside before being heated. The temper- ature of a test tube containing a solid should be raised gradually by moving it in and out of the flame, or by holding it in the flame and roll- 459 460 Experiments. ing it slightly between the thumb and forefinger. Special care must be taken to distribute the heat evenly. If the test tube contains a liquid, as is usually the case, only that part containing the liquid should be heated ; the test tube should also be inclined so that the greatest heat is not directed upon the thin bottom. When the liquid begins to boil, the test tube should be removed from the flame for an instant or held over it. In some experiments test tubes can be held be- tween the thumb and forefinger without discomfort. If they are too hot to handle, a test-tube holder may be used (Fig. 88). 3. Cutting and bending Glass Tubing. (a) Cut- ting. Determine the length needed, lay the tube on the desk, and with a forward stroke of a triangular file make a short but deep scratch where the tube is to be cut. Grasp the tube in both hands, and hold the thumbs together behind the scratch. Now push gently with the thumbs, pull at the same time with the hands, and the tube will break at the desired point. The sharp ends should be smoothed by rubbing them with emery paper or by rotating them slowly in the Bunsen flame until a yellow color is distinctly seen or until the ends become red-hot. () Bending. Glass tubes are bent in a flat flame. An ordinary illuminating gas flame may be used, but the Bunsen flame can be flattened by a wing-top attachment (Fig. 89), which slips over the top of the burner tube. The flattened Bunsen flame should be slightly yellow and about 7 centimeters (2.5 inches) wide for ordinary bends. A right-angle bend is made as follows: Determine the point at which the tube is to be bent. Grasp the tube in both hands, and hold it so that the part to FIG. 88. Test tube and holder. FIG. 89. Wing- top attachment for Bunsen burner. be bent is directly over the FIG. 90. Bending a tube into a right angle I. Introduction. 461 flame. Slowly rotate it between the thumbs and forefingers, and gradually lower it into the position shown in Figure 90. Continue to rotate it until the glass feels soft and ready to yield. Then remove it from the flame, and slowly bend it into a right angle, as shown in Figure 91. It is con- venient to have at hand a block of wood or some other right- angled object to assist the eye in completing the bend into an . ,. , Tr FIG. 91. Bending a tube into a right exact right angle. If a Bunsen angle II. flame is used, the bent part of the tube should be annealed, i.e. cooled slowly. This is done by holding it in a yellow flame until it becomes coated with soot. It should then be placed on a block of wood, and when cold wiped clean. Tubes can be bent into an oblique angle by heating them through about twice the space required for a right angle ; a very slight bend, however, is often made by holding the tube across the flame and heating a short space. Glass tubes which have been correctly bent never have flattened curves ; nor are they twisted, i.e. all parts lie in the same plane. (c) Drawing. Glass tubes can be drawn to a finer bore or into two pointed tubes as follows : Heat the glass as in (b) through about 2.5 centimeters (i inch) of its length, remove from the flame and slowly pull it apart a short distance ; let it cool for a few seconds, and then pull it quickly to the desired length. The operation is well illustrated by making a glass stirring rod. Select a piece of rod about 25 centimeters (10 inches) long and .5 FIG. 92. Stirring rods ready to be cut. centimeter (^ inch) in diameter. Heat it in the middle in the ordinary not flat Bunsen flame, and when soft draw it out slowly into the shape shown in Figure 92. Cut it into two rods by making a slight scratch where the dotted line indicates. Round off the rough edges by heating them slightly in the flame. 462 Experiments. 4. Filtering. A solid may be separated from a liquid by filtering. A circular piece of porous paper is folded to fit a glass funnel, and when the mixture is poured upon this paper, the solid the residue or precipi- tate is retained, while the liquid the filtrate passes through and may be caught in a test tube or any other vessel. The filter paper is prepared for the funnel by folding it successively into the shapes shown in Figures 93 and 94, and then opening the folded paper so that three thicknesses are on one side and one on the other (Fig. 95). The cone-shaped paper is next placed in the funnel and wet with water, FIG. 93. Folded filter paper I. FIG. 94. Folded filter paper II. FIG. 95. Folded filter paper ready for funnel. so that it will stick to the sides of the funnel and filter rapidly. The paper should never extend above the edges of the funnel, but its apex should always project slightly into the stem. The liquid to be filtered should be poured down a glass rod which touches the edge of the test tube ; the lower end of the rod should just touch the paper inside the funnel, so that the liquid will run down the side and thereby avoid bursting the apex of the filter paper. It is also advisable to adjust the apparatus so that the end of the stem of the funnel rests against the side of the vessel catching the filtrate. A funnel can be supported by standing it in a test tube, a bottle, or the ring of an iron stand. 5. Constructing and arranging Apparatus. The various parts of an apparatus should be collected, prepared, and put together be- fore starting the experiment in which the apparatus as a whole is used. The different parts which are to fit each other should be selected and arranged so that all joints are gas-tight, and as a final precaution the apparatus should be tested for leaks. All leaks should be stopped up before the apparatus is used. The following hints will be helpful : FIG. 96. Rubber tube cut at an angle. (1) To insert a glass tube into rubber tubing. Cut the rubber tubing at an angle, as shown in Figure 96, moisten the smoothed end of the glass Introduction. 463 tube with water, place the end of the glass tube in the angular-shaped cavity so that both tubes are at about a right angle, and then slip the rubber tube slowly up and over the end of the glass tube. If the glass tube is large or the rubber stiff, the rubber tube must be held firmly between the thumb and forefinger to keep it from slipping off until it is securely adjusted. (2) To fit a glass tube to a stopper. Moisten the end with water and grasp the tube firmly about 3 centimeters (i inch) from the end; hold the stopper between the thumb and 'forefinger of the other hand, and work the tube into the hole by a gradual rotary motion. Proceed in the same manner if the tube is to be pushed through the stopper. Never point the tube toward the palm of the hand which holds the stopper. Never grasp a safety tube or any bent tube at the bend when inserting it into a stopper it may break and cut the hand severely. (3) To bore a hole in a cork. Rubber stoppers are preferable, but if corks are used, they can be bored as follows : Select a cork free from cracks or channels and use a borer which is one size smaller than the desired hole. Hold the cork between the thumb and forefinger, press the larger end against a firm but soft board, and slowly push the borer by a rotary movement through the cork, taking care to keep the borer perpendicular to the cork. If the hole is too small, enlarge it with a round file. If corks are used instead of rubber stoppers, the apparatus should always be tested before use by blowing into it, stopping of course all legitimate outlets. A poor cork often means a failure, to say noth- ing of wasted time. (4) To make a platinum test wire. Rotate one end of a piece of glass rod, about 10 centimeters (4 inches) long, in the flame until it softens. At the same time grasp a piece of platinum wire about 7 cen- timeters (3 inches) long firmly in the forceps about i centimeter (.5 inch) from the end, and hold it in the flame. When the rod is soft enough, gently push the hot wire into the rod. Cool the rod gradually FIG. 97. Platinum test wire. by rotating it in the flame. The completed wire is shown in Figure 97. If a glass tube is used instead of a rod, it should be drawn out to a very small diameter (see 3 (0) before inserting the platinum wire, but in other respects the two operations are practically identical. 464 Experiments. 6. Manipulation. Ability to use apparatus rapidly, accurately, and neatly is acquired only by experience, but the following suggestions will facilitate the acquisition of this needful skill : (i) Pouring liquids and transferring solids, (a) Liquids can be poured from a vessel without spilling, by moistening a glass rod with the liquid and then pouring it down the rod as is shown in Figure 98. The angle at which the rod is held varies with circumstances. This is a convenient way to FIG. 98. Pouring a liquid down a glass rod. i- -j r i pour a liquid from a vessel containing a solid without disturbing the solid. () Liquids can often be poured from a bottle by holding the bottle as shown in Figure 99. Notice that the stopper and bottle are held in the same hand. This is ac- complished by holding the palm of the hand upward and removing the stopper by grasping it between the fingers before the bottle is lifted. All stoppers should be removed this way when possible, and not laid down, because the impurities ad- hering to the stopper may run down into the bottle and contaminate the solu- tion. The drop on the lip of the bottle should be touched with the stopper before the latter is put into the bottle ; this simple operation prevents the drop from running down the outside of the bottle upon the label or upon the shelf. (<:) Solids should never be poured directly from a large bottle into a test tube, retort, or similar vessel. A convenient method is as follows : Rotate FIG. 99. The way in which a glass stopper should be held while a liquid is being poured from a bottle. FlG. loo. Pouring a solid into a vessel with a small opening. the bottle slowly so that the Introduction. 465 solid will roll out in small quantities ; catch the solid on a narrow strip of paper folded lengthwise, and slide the solid from the paper into the desired vessel. The last part of the operation is shown in Figure 100. (2) Collecting gases. Gases are usually collected over water by means of a pneumatic trough, a common form of which is shown in Figure 102. The vessel to be filled with gas is first filled with water, covered with a piece of filter paper, inverted, and placed mouth downward on the shelf of the trough, which is previously filled with water just above the shelf. The paper is then removed, and the vessel slipped over the hole in the shelf of the trough. Glass plates instead of filter paper may be used to cover the bottle. The gas which is evolved in the generator passes through the delivery tube, and bubbles up through the water into the bottle, forcing the water out of the bottle as it rises. All gases insoluble in water are thus collected. Some heavy gases, such as hydrochloric acid, chlorine, and sulphur dioxide, are collected by allowing the gas to flow downward into an empty bottle, and displace the air in the bottle, i.e. by downward displacement. Ammonia and other light gases are usually collected by allowing the gas to flow upward into a bottle, i.e. by upward displacement. (3) Weighing and measuring. These operations are best learned by personal direction from the teacher, together with patient application of a few general principles. The following hints, however, will be of assistance : (a) Learn as soon as possible how to use the scales and interpret the weights. (b) Always leave the scales and weights in a clean, usable condition. (c) Substances should not be weighed on the bare scale pan, but on a smooth piece of paper creased on the edges or along the middle. Take the solid from the bottle with a clean spoon or spatula or --- pour by rotating the bottle as described in 6 (c). In many experiments only ap- -" 12- proximate quantities are needed. If you weigh out too much, do not put it back into the bottle, but throw it away or put it into a special bottle. (d) Liquids are measured in graduated FlG . IOI .^ Meniscus. Correct cylinders. The lowest point of the curved reading is along line I. 466 Experiments. surface of the liquid is its correct height (see Fig. 101). The average ordinary test tube holds about 30 cubic centimeters, while the large test tube so often mentioned in the succeeding experiments holds about 75 cubic centimeters. Time can be saved by remembering these volumes. (>) All measurements in this book are in the metric system (see App. i). The common denominations, their abbreviations, and English equivalents should be learned. 7. Smelling and Tasting. Unfamiliar substances should never be tasted or smelled except according to directions, and even then with the utmost caution. Never inhale a gas vigorously, but waft it gently with the hand toward the nose. Taste acids, etc., by touching a minute portion to the tip of the tongue, and as soon as the sensation is detected, reject the solution at once never swallow it. EXPERIMENTS. PHYSICAL AND CHEMICAL CHANGES. Experiment 1. Physical Change. Materials: Sugar, glass rod. Dissolve a little sugar in a test tube one fourth full of water. Dip a glass rod into the liquid and taste it. Has the characteristic property of the sugar been changed ? Dip the rod into the liquid again, and hold it over the flame of the Bunsen burner. 'As the water evaporates, a white solid appears. Taste it. What is it ? Have its original proper- ties been destroyed ? What kind of a change did they undergo ? What kind of a change did the sugar undergo ? What caused the change ? Experiment 2. Physical Change. Material: Iodine. Drop a small crystal of iodine into a dry test tube, and gently heat the bottom. As the violet vapor arises, remove the tube from the flame and let it cool. Do the crystals which form near the top resemble the original crystal ? When gently heated, do they change into the violet vapor ? How has the iodine crystal been changed ? What caused the change ? Do the original properties reappear after the cause has been removed ? What kind of a change has the iodine undergone ? Experiment 3. Physical Change. Material: Glass rod. Rub a glass rod briskly on a piece of cloth, and hold it near small bits of dry paper. Describe what happens. After a moment touch the paper again. Is the result the same ? Try again. Are the original properties of the glass restored when the cause of its change is re- moved ? What kind of a change did the glass undergo ? Experiment 4. Chemical Change. Materials : Copper wire, dilute nitric acid, iron nail, forceps. (a) Examine a piece of copper wire and notice especially its color. Grasp one end of the wire with the forceps, and hold the other end in 1 467 468 Experiments. the flame until a definite change occurs. Then remove it from the flame, and examine. Has it been changed ? Do the original properties of the copper reappear when the heated wire is cool ? What kind of a change has the copper undergone ? Has the change produced another substance ? (^) Slip another piece of copper wire into a test tube one fourth full of dilute nitric acid. Notice any change. Warm the liquid gently, and notice any additional change. What are the evidences of chemical change ? What caused the change ? What assisted or hastened it ? How has the copper been changed ? (Save the test tube and contents for (,).) (V) Carefully slip an iron nail into the liquid remaining from ($) ; let it stand a short time. Then remove and examine the coating. How does it compare with the original copper used in (a) ? What kind of a change occurred ? What caused it ? ANSWER : (1) What are the evidences of chemical changes in this experiment? (2) If a known weight of copper had been consumed in (), could it have been obtained without loss in (c) ? (3) Did the changes in this experiment involve any loss of copper? (4) What is the evidence that new substances were produced in (#) and () ? (5) What physical changes occurred in (#) and (fr) ? OXYGEN. Experiment^. Preparation of Oxygen. Materials: jjjjgrarns potassium chlorate, 1 5 grams manganese dioxide, g^bottles (about 250 cubic centimeters each), filter paper, thin piece of soft wood, sulphur, deflagrating spoon, piece of charcoal fastened to a wire, piece (about 1 5 centimeters or 6 inches) of wire picture cord unwound at one end. The apparatus is shown in Figure 102. A is a large test tube provided with a one-hole rubber stopper, to which is fitted a short glass tube, B ; the delivery tube. I), is attached to the short glass tube by the rubber tube, C. (Directions for constructing and arranging the apparatus may be found in the Introduction, 5.) Weigh the potassium chlorate on a piece of paper creased lengthwise, and slip it into the test tube ; do the same with the manganese dioxide. Shake the test tube until the chemicals are thoroughly mixed ; then hold Oxygen. 469 the test tube in a horizontal position and roll or shake it until the mix- ture is spread along the tube its entire length. Insert the stopper with its tubes, and clamp the test tube to the iron stand, as shown in the FIG. 102. Apparatus arranged for preparing oxygen. figure, taking care not to crush the tube ; the test tube should incline toward the trough, to prevent any water from flowing back upon the hot glass. Fill the pneumatic trough with water until the shelf is just covered. Fill the bottles/)/// of water, cover each with a piece of filter paper, in- vert them in the trough, and remove the filter paper ; leave two bottles on the shelf and three on the bottom. The end of the delivery tube should rest on the bottom of the trough, just under the hole in the shelf. Heat the whole test tube gently with a flame about 8 centimeters (or 3 inches) high. When the gas bubbles regularly through the water, slip a bottle over the hole. The gas will rise in the bottle and force out the water. Move the flame slowly along the test tube, but concentrate the heat toward the closed end, and always keep the flame behind any water which may be driven out of the mixture. If the gas is evolved too rapidly, lessen the heat; if too slowly, increase it ; if not at all, examine the stopper and the rubber connecting tube for leaks, and adjust accordingly. When the first bottle of gas is full, remove and cover it with a piece of wet filter paper, and slip another bottle over the hole. When five bot- tles of gas have been collected, remove the end of the delivery tube from the water, lest the cold water be drawn up into the hot test tube as the gas contracts. Perform the next experiment at once. 470 Experiments. Experiment 6. Properties of Oxygen. Proceed as follows with the oxygen prepared in the preceding experiment. (a) Dip a glowin^_stickjDf^wood into one bottle, and observe the change. Remove the stick, and repeat as many times as possible. Does the gas burn? How does the glowing stick change? What property of oxygen does this experiment show? (b} Put a small piece of sulphur in the deflagrating spoon, hold the spoon in the flame until the blue flame of the burning sulphur can be seen, then lower the spoon into a bottle of oxygen. Notice the change in the flame. Describe it. Brush a little of the vapor cautiously toward the nose. Of what does the odor remind you? (Plunge the spoon into water to extinguish the burning sulphur, and covef the bottle with* a piece of filter paper.) (c) Hold the charcoal in the flame long enough to produce a faint glow, then lower IFmto a bottle of oxygen. Describe the result. (a) Melt the sulphur in the deflagrating spoon, and dip the unwound end of the wire picture cord into the melted sulphur. Lower the end coated with burning sulphur into a bottle of oxygen. The iron wire should burn brilliantly. Describe the change. Sometimes the sub- stance produced by the change coats the inside of the bottle Describe it, if it is visible. (tf) With the remaining bottle, repeat any of the above experiments. EXERCISES : (1) Write a brief account of the above experiments in your note book, answering all questions and directions. (2) Sketch the apparatus used to prepare oxygen. (3) Summarize the properties of oxygen. (4) What is its most characteristic property? (NOTE. The test tube used in Experiment 5 may be cleaned with warm water.) Experiment 7. Preparation of Oxygen from Mercuric Oxide. Materials : Mercuric oxide, stick of wood. Put a little mercuric oxide on the end of a narrow piece of paper creased lengthwise, and slip the powder into a test tube. The pow- der should nearly fill the round end of the test tube. Hold the test tube in a horizontal position, shake it to spread the powder into a thin Hydrogen. 471 layer, attach the test-tube holder, and heat the test tube (still horizontal) in the upper part of the Bunsen flame. Do not heat one place, but move the tube back and forth. As soon as a definite change is noticed inside the tube, insert a glowing stick of wood. Observe and describe the change. If there is no change, heat strongly, and test again. .What gas is liberated? Observe the deposit inside the tube. What is it? If its nature is doubtful, let the tube cool, and examine again. EXERCISES : (1) Describe briefly the whole experiment. (2) What historical interest has this experiment? (NOTE. If the test tube has been partially melted, save it for a sub- sequent experiment.) HYDROGEN. Experiment 8. Preparation of Hydrogen. Materials : Granu- lated zinc, dilute sulphuric acid, pneumatic trough, four bottles, filter paper, taper, matches. The apparatus is shown in Figure 103. A is a large test tube provided with a two-hole stopper, through which passes the safety tube, B, and the right-angle bend, C\ the long (15 cm. or 6 in.) delivery tube, E, is attached to the bent tube by the rubber tube, D. Precaution. Keep all flames away from the hydrogen generator. Fill the test tube half full of granulated zinc as follows : Crease a piece of paper lengthwise, pour the zinc from the bottle upon the paper, incline the test tube, and slip the zinc .into it from the paper do not drop it in. Insert the stopper with its tubes ; if the end of the safety tube does not go in easily, hold the test tube in a horizontal position and shake the zinc about, and at the same time push the stopper gently but firmly into place. Clamp the apparatus into the position shown in the figure or stand it -in a test-tube rack. Fill the pneumatic trough with water as before, and ad- just the apparatus so that the end of the delivery tube rests on the bottom of the trough FIG. 103.- Apparatus for preparing hydrogen. under ^ ho]e ^ ^ ^^ Fill the bottles with water and invert them in the trough, as in Experiment 5. O 472 Experiments. Pour enough dilute sulphuric acid through the safety tube to fill the test tube about half full, taking care to leave a little acid in the lower bend of the safety tube. This precaution prevents the gas from escap- ing from the back of the apparatus ; if at any time the gas should flow backward, pour a little acid into the bend ; if the acid does not flow down the safety tube, loosen the stopper for an instant. As soon as the interaction of the zinc and sulphuric acid produces hydrogen, the gas will bubble freely through the water in the trough. Slip a bottle over the hole, and collect and remove the bottle of gas as in Experi- ment 5, taking care to cover the bottle firmly with a piece of wet filter paper. If the evolution of gas slackens or ceases, add a little more acid through the safety tube. Collect four bottles of hydrogen, and proceed at once with the next experiment. Experiment 9. Properties of Hydrogen. Study as follows the hydrogen gas prepared above : (a} Uncover a bottle for an instant to let a little air in, and then drop a lighted match into the bottle. Describe the result. () Remove the paper from a bottle of hydrogen, and allow it to remain uncovered for three minutes by the clock. Then show the presence or absence of hydrogen by dropping a lighted match into the bottle. Describe the result. What property of hydrogen is shown by this experiment? (c} Verify your answer to the last question, thus : Hold a bottle of air over a covered bottle of hydrogen, remove the paper, and bring the mouths of the bottles close together. (See Fig. i.) Hold them there for a minute or two, then stand the bottles on the desk and cover them with wet filter paper. Drop a lighted match into each bottle. What has become of the hydrogen? How does (c) verify ()? (d} Invert a covered bottle of hydrogen, remove the paper, and quickly thrust a lighted taper up into the bottle. Withdraw the taper and then insert it again. Does the hydrogen burn? If so, where? Does the taper burn when in the bottle? When out of the bottle ? Feel of the neck of the bottle ; describe and explain. What three properties of hydrogen are shown by this experiment ? Experiment 10. Burning Hydrogen. (Teacher's Experi- ment.) Materials: Apparatus shown in Figure 2, which consists of a Hydrogen. 473 500 cubic centimeter flask fitted with a two-hole rubber stopper, safety tube, and double right-angle bend ; the last is attached to a U-tube, which is also connected to a delivery tube provided with a short piece of capillary glass tubing; calcium chloride, small bottle, platinum wire, cotton, granulated zinc, dilute sulphuric acid. Fill the U-tube two thirds full of calcium chloride, put a wad of cotton beneath the stopper of each arm, and connect the U-tube with the generator and the delivery tube. Stand the apparatus on the table, examine all joints to be sure they are tight, extinguish all flames in the vicinity, and proceed exactly according to the following directions : Pour slowly but continuously through the safety tube enough (about 50 cubic centimeters) dilute sulphuric acid upon at least 25 grams of granulated zinc to produce a steady current of hydrogen gas for about five minutes. It is advisable to use considerable zinc and a moderate amount of acid. Acid must not be added after the evolution of gas begins, unless, of course, the experiment is begun anew. Let the gas bubble through the acid for at least two minutes by actual observation, then attach the capillary tube by the rubber connector to the end of the delivery tube, leaving a short space between the ends of the two glass tubes so that the rubber tube may be compressed suddenly, if necessary. Let the gas run for another full minute. This latter pre- caution is to drive all air out of the capillary tube. Light the hydrogen, and observe at once the nature of the flame, its color, heat (by holding a match or platinum wire over it), and any other striking property. Then hold a small dry bottle over the flame in such a position that the flame is just inside the bottle. When conclusive evidence of the prod- uct of burning hydrogen is seen inside the bottle, remove the bottle, and extinguish the flame at once by pinching the rubber connector. Remove the generator to the hood, and if the evolution of hydrogen is still brisk, dilute the acid by pouring water through the safety tube. Examine the inside of the bottle. What is the deposit ? Explain its formation. EXERCISES FOR THE CLASS: (1) What does this experiment suggest about the composition of water ? (2) Does this experiment illustrate oxidation? Why? Synthesis? Why? (3) Describe the whole experiment, and sketch the apparatus. 474 Experiments. WATER. Experiment 11. General Distribution of Water. Materials: Wood, meat, potato. Heat successively in dry test tubes a small piece of wood, of meat, or of potato (or any other fresh vegetable) . Hold the open end of the test tube lower than the other end. Is there conclusive evidence of water? Since most animal and vegetable substances act similarly, what general conclusion can be drawn from this experiment ? Experiment 12. Simple Tests for Impurities in "Water. Materials: Distilled water, water containing dirt, a sulphate, a chloride, and a lime compound ; nitric acid, ammonium hydroxide, acetic acid, sulphuric acid (concentrated), solutions of potassium permanganate, silver nitrate, barium chloride, ammonium oxalate ; and limewater. (a) Organic Matter. Fill a clean test tube half full of distilled water, and another with water containing a little dirt or a bit of paper. Add to each test tube a drop or two of concentrated sulphuric acid and suffi- cient potassium permanganate solution (made from distilled water) to color each liquid a light purple, as nearly alike as possible. Label one tube, and then heat gently nearly to the boiling point the tube contain- ing the impure water. As soon as a definite change is seen, heat the other cautiously, as too sudden heat may cause the liquid to "bump out." Organic matter decolorizes potassium permanganate solution. Which tube shows the more organic matter? (b) Chlorides. To a test tube half full of distilled water add a few drops of nitric acid, and then a few drops of silver nitrate solution. Do the same with water known to contain a chloride in solution. What is the difference between the results ? The cloudiness, or solid, is due to silver chloride, which is always formed when silver nitrate is added to hydrochloric acid or a chloride in solution (chlorides being closely related to hydrochloric acid). Silver chloride is soluble in ammonium hydroxide. Try it. This is the usual test for chlorides (and conversely for soluble silver compounds), and will hereafter be used without further description. (c} Sulphates. To a test tube half full of distilled water add a few drops of sulphuric acid and a few drops of barium chloride solution. The white precipitate is barium sulphate. It is insoluble in all common liquids, and is always formed when barium chloride is added to sulphu- ric acid or a sulphate in solution (sulphates being closely related to sul- phuric acid). Test the impure water for sulphates. Water. 475 (y ) Lime Compounds. Add a few drops of a fresh solution of ammo- nium oxalate to a test tube half full of clear limewater. Limewater is a solution of calcium hydroxide, and the white precipitate formed is calcium oxalate, which is soluble in hydrochloric acid but not in acetic acid. Try it. This is the test for calcium compounds, often called "lime" compounds, because lime, which is calcium oxide, is so well known. Apply this test to distilled water and to water known to con- tain calcium compounds, and compare the two results. (e) Summarize briefly the whole experiment. (NOTE. If time permits, this experiment should be applied by the class to water whose impurities are unknown.) Experiment 13. Distillation. (Teacher's Experiment.) Ma- terials: Condenser, etc., shown in Figure 6, potassium permanganate, impure water, and solutions used in Experiment 12. Fill the flask, C, half full of water known to contain the impurities mentioned in Experiment 12, add a few crystals (3 or 4) of potassium permanganate, and connect with the condenser as shown in Figure 6. Attach the inlet tube to the faucet, fill the condenser slowly, and regu- late the current so that a small stream flows continuously from the outlet tube into the sink or waste pipe. Heat the liquid in C gradually, and when it boils, regulate the heat so that the boiling is not too vio- lent. As the distillate collects in the receiver, Z?, test separate portions for organic matter, chlorides, sulphates, and calcium compounds. EXERCISES FOR THE CLASS : (1) Is organic matter found ? (2) Is mineral matter found ? (3) If the distilling liquid had contained a volatile substance, like ammonia or alcohol, would the distillate contain such a substance ? Experiment 14. Solubility of Gases. (a} Warm a little faucet water in a test tube. Is there immediate evidence of a previously dissolved gas ? Is there evidence of much gas ? What effect has increased heat ? (6) Warm slightly a few cubic centimeters of ammonium hydroxide in a test tube. Do the results resemble those in (a) ? As soon as the final result is obtained, pour the remaining liquid down the sink and flush well with water. 476 Experiments. (V) Repeat (), using a little concentrated hydrochloric acid. Do the results resemble those of (a) and (b) ? ANSWER : (1) How does increased temperature affect the solubility of gases ? (2) What gases dissolve freely in water ? Experiment 15. Solubility of Liquids. Materials: Alcohol, kerosene, glycerine, carbon disulphide. (a) To a test tube half full of water add a little alcohol and shake. Is there evidence of solution ? Add a little more and shake. Add a third portion. Is there still evidence of solution ? Draw a conclusion as to the solubility of alcohol in water. () Repeat (), using successively kerosene, glycerine, and carbon disulphide. Observe the results and conclude accordingly. (c) Summarize the results in a table. Experiment 16. Solubility of Solids. Materials: About 20 grams of powdered copper sulphate, 6 grams of powdered potassium chlorate, i gram of calcium sulphate. (a) Label three test tubes I, II, III. Fill each about one third full. To I add i gram of powdered copper sulphate, to II add i gram of powdered potassium chlorate, to III add i gram of calcium sulphate. Shake each test tube, and then allow them to stand undisturbed for a few minutes. Is there evidence of solubility in each case? Is there evidence of a varying degree of solubility? If III is doubtful, carefully transfer a portion of the clear liquid to an evaporating dish by pouring it down a glass rod (see Introduction, 6 (i )(#)), and evaporate to dry- ness. Is there now conclusive evidence of solution? Draw a general conclusion from this experiment. Save solutions I and II for (). Tabulate the results of (d) as follows, using the customary terms to express the degree of solubility : TABLE OF SOLUBILITY OF TYPICAL SOLIDS. SOLUTE. SOLVENT. RESULTS. i . Copper sulphate 2. Potassium chlorate Water at tempera- ture of labora- I. 2. 3. Calcium sulphate tory. 3- Water. 477 () Heat I and add gradually 4 more grams of powdered copper sulphate. Does it all dissolve? Heat II and add 4 more grams of powdered potassium chlorate. Does it all, or most all, dissolve? What general effect has increased heat on the solubility of solids? What is the difference between this general result and that in Experiment 14? Save the solutions for (c) . (c) Heat I and II nearly to boiling, and as the temperature in- creases add the respective solids. Do not boil the liquid away. Is there a limit to their solubility? Draw a general conclusion from these typical results. Experiment 17. Supersaturation. Material: Sodium thio- sulphate. Fill a test tube nearly full of crystallized sodium thiosulphate and add a very little water. Warm slowly. As solution occurs, heat gradually to boiling. Add sodium thiosulphate until no more will dissolve. Pour the solution into a warm, clean, dry test tube and let it stand until cool. Then drop in a small crystal of sodium thio- sulphate and watch for any simple but definite change. What hap- pens? Is the excess of solid large? How does a supersaturated solution differ from a saturated one? Experiment 18. Water of Crystallization. Materials: Crys- tallized sodium carbonate, gypsum, copper sulphate, evaporating dish, gauze-covered ring (or tripod). (a) Heat a few small crystals of sodium carbonate in a dry test tube, inclining the test tube so that the open end is the lower. What is the evidence that they contained water of crystallization? If there is any marked change in the appearance of the crystals, describe and explain it. (b) Repeat, using a crystal of gypsum. Answer the question asked in (a). (c} Heat two or three small crystals of copper sulphate in an evapo- rating dish which stands on a gauze-covered ring. As the action pro- ceeds, hold a dry funnel or glass plate over the dish. Is there conclusive evidence of escaping water of crystallization ? Do the crystals change in color? In shape? Can the form of the crystals be changed by gently touching the mass with a glass rod? Continue to heat until the resulting mass is a bluish gray. Let the dish cool. Meanwhile heat a test tube one half full of water. When the dish has cooled somewhat, 478 Experiments. pour the hot water slowly into the dish upon the copper sulphate. Ex- plain the change in color, if any. If there are any lumps, crush them with a glass rod. Let the clear solution evaporate for several hours. Are crystals deposited? If not, heat' a few minutes, and cool again. If so, why ? Have they water of crystallization, and, if so, where did they get it? Experiment 19. Efflorescence. Put a fresh crystal of sodium carbonate and of sodium sulphate on a piece of filter paper, and leave them exposed to the air for an hour or more. Describe any marked change. What does this change show about the air ? About the crystal ? Experiment 20. Deliquescence. Put on a glass plate or block of wood a small piece of granulated calcium chloride and of sodium hydroxide. Leave them exposed to the air for an hour or more. Describe any marked change which takes place. Compare the action with that of Experiment 19. Experiment 21. Solution and Chemical Action. Materials: Powdered tartaric acid, sodium bicarbonate, lead nitrate, potassium di- chromate, mortar, dish of water. (#) Mix in a dry mortar small but equal amounts of powdered tar- taric acid and sodium bicarbonate. Is there any decided evidence of chemical action ? Pour the mixture into a dish of water. Is there con- clusive evidence of chemical action ? (b) Repeat, using powdered lead nitrate and powdered potassium dichromate. Describe the results in (a) and (b). How does solution influence chemical action ? Why are so many solutions used in the laboratory ? Experiment 22. Electrolysis of Water. (Teacher's Experi- ment.) Materials : Hofmann apparatus, sulphuric acid, taper, matches, short piece of capillary glass tubing. Fill the Hofmann apparatus, Figure 10, with water containing 10 per cent of sulphuric acid, so that the water in the reservoir tube stands a short distance above the gas tubes after the stopcock in each has been closed. Connect the platinum terminal wires with a battery of at least two cells. As the action proceeds, small bubbles of gas rise and collect Water. 479 at the top of each tube. Allow the current to operate until the smaller volume of gas is from 8 to 10 centimeters in height. Measure the height of each gas column. Assuming that the tubes have the same diameter, the volumes are in approximately the same ratio as their heights. How do the volumes compare ? Test the gases as follows : (#) Hold a glowing taper over the tube containing the smaller quantity of gas, cautiously open the stopcock to allow the water (or air) to run out of the glass tip, and then let out a little gas upon the glowing taper. What is the gas ? Repeat until the gas is exhausted. Care must be taken not to lose the gas. It is ad- visable to have at hand several partially burned tapers or thin splints, in case any escaping water extinguishes the first one. (t>) Open the other stopcock long enough to force out the water in the glass tip ; close the stopcock, and, by means of a short rubber tube, attach the capillary tube close to the end of the glass tip. Open the stopcock again, let out the gas slowly, and hold at the same time a lighted match at the end of the tip, then immediately thrust a taper into the small and almost colorless flame. What is the gas ? Repeat until the gas is exhausted. EXERCISES FOR THE CLASS: (1) Describe the whole experiment. (2) Draw a general conclusion from this experiment. (3) What does this experiment show about the composition of water ? (4) Sketch the apparatus. Experiment 23. Interaction of Water and Chlorine. (Teach- er's Experiment.) Materials: Glass\ube I meter long and about 2 centimeters in diameter, cork for one end, evaporating dish, chlorine water. Construct a chlorine generator, as described in Experiment 38, and prepare about 250 cubic centimeters of chlorine water by causing the gas to bubble through a bottle of water until the water smells strongly of the gas. Close one end of the tube with a cork. The cork must fit air tight, and as a precaution should be smeared (after insertion) with vaseline or coated with paraffin. Fill the tube full of chlorine water, cover the open end with the thumb or finger, invert the tube, and immerse the open end in the evaporating dish, which should be nearly 480 Experiments. full of chlorine water. Clamp the tube in an upright position, and stand the whole apparatus where it will receive the direct sunlight for at least six hours. Bubbles of gas will soon appear, rise, and collect at the top. When sufficient gas for a test has collected, unclamp the tube, cover the open end with the thumb or finger, invert the tube, and put a glowing taper into the gas. Repeat as long as any of the gas remains. EXERCISES FOR THE CLASS: (1) What gas is produced by the interaction of chlorine and water ? (2) Describe this experiment. (3) What does it show about the composition of water ? (4) Sketch the apparatus. Experiment 24. Interaction of Water and Sodium. Mate- rials : Sodium, pneumatic trough filled with water as usual, tea lead, for- ceps, red litmus paper. Precaution. Sodium, shotdd be handled cautiously and used strictly according to directions. Small fragments must not be left about nor thrown into the refuse jar, but into a large vessel of water especially pro- vided for that purpose. (a) If the sodium is brown, scrape off the coating. Cut off a piece of sodium not larger than a small pea, and drop it upon the water in the trough. Stand far enough away so that you can just see the action. Wait until you are sure the action has stopped, and then describe all you have seen. (b) The action in (a) may be further studied as follows : Fill a test tube with water, invert it, and clamp it in the trough so that the mouth is over the hole in the shelf of the trough. Wrap a small piece of sodium loosely in a piece of tea lead about 5 centimeters (2 inches) square, make two or three small holes in the tea lead, and then thrust it under the shelf of the trough with the forceps. A gas will rise into the test tube. Proceed similarly with additional small pieces of sodium and dry tea lead until the test tube is nearly full of gas ; then unclamp and remove, still keeping the tube inverted. Hold a lighted match, for an instant, at the mouth of the tube. Observe the result, watching especially the mouth of the tube. What is the gas? Why? Remembering that sodium is an element, where must the gas have come from? If there is any doubt about the nature of the gas, collect more, and subject it to those tests which will prove its nature. The Air. 481 (V) Put a piece of filter paper on the water in the trough, and before it sinks drop a small piece of sodium upon it. Stand back and observe the result. Wait for the slight explosion which usually occurs soon after the action stops. Describe all you have seen. What burned? What caused it to burn? To what is the vivid color probably due? (In answering these questions, utilize your knowledge (i) of the prop- erties of the gases previously studied, and (2) of the usual accompani- ment of chemical action, suggested here by the melting of the sodium.) (d} Test the water in the trough with red litmus paper. Push the paper to the bottom or to the place where it is certain that chemical action between water and sodium has taken place. Test until the red litmus paper has undergone a decided change in color. Describe this final result. With another piece of red litmus paper test a solution of sodium hydroxide. Is the result similar? Dip a glass rod or the plati- num test wire (see Int. 5 (4)) into this solution and hold it in the Bunsen flame. Describe the result. Is the color of this flame and that noticed in (<:) the same? Are the dissolved substances identical? (e) W T hat does the whole experiment show about the composition of water ? THE AIR. Experiment 25. Composition of the Air. Materials: Solu- tions of pyrogallic acid and potassium hydroxide, 1 pneumatic trough half filled with water at the temperature of the room, 500 and 25 cubic centi- meter graduated cylinders. The apparatus consists of an Erlenmeyer flask (250 cubic centimeters) provided with a one-hole rubber stopper into (but not through) which passes a short glass tube ; to the outer end of this tube, which projects 2.5 centimeters (i inch) above the stopper, a rubber tube (5 centimeters or 2 inches long) is tightly fastened ; a Hofmann screw is attached to the rubber tube close to the end of the glass tube. (a) The volume of the flask is found thus : Fill the flask completely with water from the pneumatic trough. Loosen the screw and push the stopper into the flask as far as it will go. Wipe the flask dry and carefully remove the stopper. Pour most of the water from the flask into the 500 cubic centimeter graduate, and read the volume : the last portions of the water in the flask should be poured into the 25 cubic 1 The pyrogallic acid is a 10 per cent solution, and the potassium hydroxide 50 per cent. 482 Experiments. centimeter graduate, so that the volume can be read accurately. (See Fig. 101). Record the total volume of the flask as shown in {d}. (b) Measure exactly 10 cubic centimeters of pyrogallic acid in the small graduate (see Int. 6 (3) (d)), and pour it into the flask. Add 20 cubic centimeters of potassium hydroxide solution, and insert the rubber stopper quickly and firmly. Tighten the screw. Shake the flask vigorously for a minute. Then invert it and watch the surface of the liquid for bubbles. If any appear, the apparatus leaks. Find the leak, if any, start the experiment again from (), taking care to remedy the defect before the flask is shaken. If no bubbles appear, continue to shake at intervals from fifteen to twenty minutes. During this operation the oxygen is absorbed by the solution. (c) Place the flask on its side in the water of the pneumatic trough, and open the screw, taking care (i) not to let any of the solution run out, (2) nor to let too much water run in, and (3) to keep the end of the rubber tube constantly below the surface. After the water has stopped running in, remove the flask from the trough. Open the flask, put a glowing stick into the gas, and observe the result. The gas is nitrogen. Measure carefully the volume of the final liquid in the flask. (d) Record and calculate as follows : (a) Volume of original solution = 30 cc. (b) Capacity of flask = cc. (c) Volume of air taken (b a) = , (d) Final volume of liquid = (e) Volume of water which entered (d a) (f ) Per cent of water which entered (e -s- c) But the per cent of entering water equals the per cent of gas ab- sorbed, hence (g) Per cent of oxygen (h) Per cent of nitrogen (100 g) = Experiment 26. Air contains Water Vapor. Prove by an experiment that air contains water vapor. Experiment 27. Air contains Carbon Dioxide. (a) Expose a small bottle of limewater to the air. After a short time, examine the surface of the liquid. Describe the change. Ex- plain it. Acids, Bases, and Salts. 483 () If a blast lamp (or bicycle pump) is available, replace the lamp with a glass tube, and force air through a bottle half full of limewater, until a definite change occurs. Describe it. Explain it. ACIDS, BASES, AND SALTS. Experiment 28. General Properties of Acids. Materials : Dilute sulphuric, nitric, and hydrochloric acids, glass rod, litmus paper (both colors), zinc. Fill separate test tubes one third full of each of the acids. Label the tubes in some distinguishing manner. (a) Dip a clean glass rod into each acid and cautiously taste it. Describe the taste by a single word. (b) Dip a clean glass rod into each acid and put a drop on both kinds of litmus paper. The striking change is characteristic of acids ; draw a general conclusion from it. () Slip a small piece of zinc into each test tube successively. If no chemical action results, warm gently. Test the most obvious product by holding a lighted match inside of each tube. What gas comes from the hydrochloric and sulphuric acids? (d} Summarize the general results of this experiment. Experiment 29. General Properties of Bases. Materials: Litmus paper (both colors), glass rod, sodium hydroxide and potassium hydroxide solutions, and ammonium hydroxide. (a) Rub a little of each liquid between the fingers, and describe the feeling. Cautiously taste each liquid by touching to the tip of the tongue a rod moistened in each, and describe the result. (b) Test each solution with litmus paper. Describe the result. (c) Summarize the general results of this experiment. (d) Compare acids and bases as to taste and to reaction with litmus. Experiment 30. A Property of Many Salts and All Neutral Substances. Materials: Litmus paper (both colors), glass rod, dilute solutions of sodium chloride, potassium nitrate, potassium sulphate, and barium chloride. Test each solution with litmus paper. Describe the result. Com- pare with the litmus reaction of acids and bases. Draw a general conclusion from this experiment. 484 Experiments. Experiment 31. The Nature of Common Substances. Determine by the litmus test the nature of lemon juice, vinegar, sweet and sour milk, washing soda, borax, wood ashes, faucet water, baking soda, sugar, cream of tartar, the juice of any ripe fruit and any green fruit. Make a solution of each of the solids before testing. Tabulate the results as follows : NATURE OF COMMON SUBSTANCES. ACID. ALKALINE. NEUTRAL. Experiment 32. Neutralization. Materials: Sodium hydrox- ''ide (solid), hydrochloric acid, nitric acid, silver nitrate solution, blue litmus paper, glass rod, evaporating dish, gauze-covered ring. Dissolve a small piece of sodium hydroxide in an evaporating dish half full of water. Slowly add dilute hydrochloric acid, until a drop taken from the dish upon a glass rod reddens blue litmus paper. Then evaporate to dryness by heating over a piece of wire gauze supported by a ring. Since the residue mechanically holds traces of the excess of hydrochloric acid added, it is necessary to remove this acid before applying any test. Heat the dish until all the yellow color disappears, then moisten the residue carefully with a few drops of warm water and heat again to remove the last traces of acid. This precaution is essen- tial to the success of the experiment. Test a portion of the residue with litmus paper to find whether it has acid, alkaline, or neutral properties. Taste a little. Test (a} a solu- tion of the residue for a chloride, and (b} a portion of the solid residue for sodium. (See Exps. 12 (t>) and 24.) Draw a definite conclusion from the total evidence. Heat, Light, Electricity, and Chemical Action. 485 HEAT, LIGHT, ELECTRICITY, AND CHEMICAL ACTION. Experiment 33. Heat and Chemical Action. Materials : Lime, evaporating dish, match. Put a small piece of lime in an evaporating dish, and sprinkle a little water over it. Watch for a change. If no marked change soon occurs, add a little more water. Describe the change. Touch a match to the mass. Is there evidence of much heat? What caused the heat? Experiment 34. Heat and Chemical Action. Materials: Sul- phur, powdered iron, dilute hydrochloric acid. Put about 3 grams of sulphur and 3 grams of powdered iron in a test tube. Cover the mouth of the test tube with the thumb and shake until the two substances are well mixed. Attach the test tube to the holder and heat strongly in the flame. As soon as the sulphur melts and boils and the contents give evidence of decided chemical action, remove the test tube at once from the flame, and watch the change. Is there evidence of heat? Of increasing heat? Of much heat? When the tube is cool, break the end, and examine the contents. De- scribe it. It is a compound called iron sulphide, and is the product of the chemical action which was started by heat. But the chemical action itself was so vigorous that it increased_the heat. The fact that the product differs from the original mixture may be shown as follows : Add dilute hydrochloric acid to a part of the product and also to a little of the original mixture, testing the gaseous product in each case by the odor. Is the odor the same? State briefly how heat and chemical action are related, using this experiment as an illustration. Experiment 35. Light and Chemical Action. Materials : Potassium bromide, silver nitrate solution, funnel, filter paper, glass rod. Dissolve a crystal of potassium bromide in a test tube one fourth full of water, add an equal volume of silver nitrate solution, and shake. The precipitate is silver bromide. Describe it. Filter (see Int. 4). Remove the filter paper from the funnel, unfold it, and expose the silver bromide for a few minutes to the light sunlight, if possible. Describe the change. What caused the change? How is this property of silver bromide utilized ? 486 Experiments. Experiment 36. Electricity and Chemical Action. (Teacher's Experiment.) Repeat Experiment 22. EXERCISES FOR THE CLASS: (1) Define electrolysis, electrode, electrolyte, ion, anion, cation. (2) State briefly the accepted explanation of the electrolysis of water. (3) Is hydrogen an anion or cation? At what electrode does it collect ? (4) Answer the same questions (as in 3) about oxygen. Experiment 37. Electricity and Chemical Action. (Teach- er's Experiment.) Materials : Starch, potassium iodide, mortar and pestle, filter paper, sheet tin (or iron), battery of two or more cells. Grind together in a mortar a lump of starch and a crystal of potas- sium iodide. Add enough water to make a thin liquid. Dip a strip of filter paper into the mixture, and spread the wet paper upon a sheet of tin (or iron) . Press the end of the wire attached to the zinc (of the battery) upon the tin, and draw the other wire across the sheet of paper. The marks are caused by iodine which is liberated from the potassium iodide and colors the starch. EXERCISES FOR THE CLASS: (1) Describe briefly this experiment. (2) Iodine is a non-metal. At what electrode is it liberated? Is iodine an anion or a cation ? CHLORINE. {Do not inhale chlorine?) Experiment 38. Preparation of Chlorine. Materials: Con- centrated hydrochloric acid, 30 grams manganese dioxide, bundle of fine brass wire, strip of calico, paper with writing in lead pencil and in ink, litmus paper (both colors), taper. The apparatus is shown in Fig- ure 104. It is the same as that used to prepare hydrogen ; and there are also needed four bottles, a wooden block (about 10 centimeters or 4 inches square) with a hole in the center, and four glass plates to cover the bottles. Weigh the manganese dioxide upon a piece of paper creased length- wise. Slip it into the test tube, A (see Int. 6 ()). Arrange the appa- Chlorine. 487 ratus as shown in the figure. Pour enough concentrated hydrochloric acid through the safety tube to cover the man- ganese dioxide. Heat gently with a small flame, keeping the flame below the level of the contents of the test tube. Chlorine is rapidly evolved as a greenish gas, and passes into the bottle, G, which should be removed when full (as seen by the green color) and covered with a glass plate ; the bottle may be easily removed by holding the block, F, in one hand and pulling the bottle, G, aside, bending the whole delivery tube at the same time at the rubber connection, D. If the evolution of gas slackens, add more acid through the safety tube. Collect four bottles, and perform the next experiment at once. FIG. 104. Apparatus arranged for preparing chlorine. Experiment 39. Properties of Chlorine. Study as follows the gas prepared above : (a) Heat the bundle of brass wire and thrust it into a bottle of chlo- rine. Describe the result, especially the evidence of chemical action and of new products. (b) Into a bottle of dry chlorine put a piece of calico, litmus paper (both colors), and paper contain- ing writing in black and in red ink. Allow the whole to remain undisturbed for a few minutes and then describe the change, if any. Add several drops of water, and describe the change. Draw a general con- clusion from the whole experiment. (c} Hold a burning taper in a, bottle of chlorine long enough to observe the result. Draw a conclusion. Verify it thus : Fold a strip of filter paper (about 10 centimeters or 4 inches wide) into the shape shown FIG. 105. Fluted ln Figure 105; cautiously heat 1 about 10 cubic centi- paper. meters of turpentine in a large test tube; saturate 1 Hold the test tube with the holder. Remember that turpentine ignites easily. If the turpentine catches fire, press a damp towel over it. 488 Experiments. the paper with the hot turpentine and drop it into a bottle of chlorine. Describe the result. When the action is over, examine the paper, and draw a conclusion regarding the action between hot turpentine and chlorine. Wax (in the taper) and turpentine are mainly compounds of hydro- gen and carbon. Explain the result in (c) . ANSWER : (1) Many metals act like the brass in (#). What general conclu- sion can be drawn about the reaction of chlorine and metals? (2) What is essential for the bleaching action of chlorine? (3) What does (c} show about the attraction between chlorine and hydrogen ? (4) What class of chemical changes is illustrated by (#) ? What classes by (c) ? (5) What class of chemical changes is illustrated by the preparation of chlorine? (6) What three striking properties has chlorine? How can it be distinguished from all gases previously studied? Experiment 40. Bleaching by Bleaching Powder. Mate- rials : Bleaching powder, sulphuric acid, calico. Put a little bleaching powder into a test tube and add enough water to make a thin paste. Add a few drops of dilute sulphuric acid, and then dip a strip of bright-colored calico into the mixture. Remove the calico in a few minutes, and wash it with water. Describe the change in the calico. Experiment 41. Preparation of Hydrochloric Acid. Mate- rials : The apparatus used in Experiment 38 ; 20 grams sodium chlo- ride, concentrated sulphuric acid, pneumatic trough filled with water as usual, stick of wood, litmus paper (blue), ammonium hydroxide. (a) Put 8 cubic centimeters of water in a small bottle or evaporating dish, cautiously add 12 cubic centimeters of concentrated sulphuric acid, and stir until the two are mixed. While this mixture is cooling, weigh the salt, slip it into the test tube, and then arrange the apparatus as shown in Figure 104. Pour half the cold acid mixture through the safety tube, let it settle through the salt, and then add the remaining acid. Heat gently with a low flame, as in the preparation of chlorine. Hydrochloric acid gas ' is evolved, and passes into the bottle, which Chlorine. 489 should be removed when full, as directed under chlorine. A piece of moist blue litmus paper held at the mouth of the bottle will show when it is full. Collect these bottles, cover each with a glass plate, and set aside until needed. (b) As soon as the third bottle of gas has been collected, removed, and covered, put in its place a bottle one fourth full of water. Adjust its height (if necessary) by wooden blocks so that the end of the delivery tube is just above the surface of the water. Continue to heat the generator at intervals, and the gas will be absorbed by the water. Shake the bottle occasionally. Meanwhile study the gas already collected. Experiment 42. Properties of Hydrochloric Acid Gas. Proceed as follows with the hydrochloric acid gas prepared by Experiment 41 : (#) Insert a blazing stick of wood into a bottle. Remove as soon as the change is noticed. Describe the change. Compare the action with the behavior of hydrogen and of oxygen under similar conditions. (^) Hold a piece of wet filter paper near the mouth of the same bottle. Describe the result. What is the cause? (c) Invert a bottle, and stand it upon the shelf of the pneumatic trough. Describe any change noticed inside the bottle after a few minutes. What property of the gas does the result illustrate ? Verify the observation by a simple test applied to the contents of the bottle. (d) Drop into the remaining bottle of gas a piece of filter paper wet with ammonium hydroxide. Describe the result. What name has the product? (e) State other properties of hydrochloric acid gas which you have observed ; e.g. color, odor, density. Proceed at once with the next experiment. Experiment 43. Properties of Hydrochloric Acid. Remove the bottle in which the hydrochloric acid gas is being ab- sorbed (see Exp. 41 ()), and study the solution as follows : (a) Determine its general properties, e.g. taste (cautiously), action with litmus, and with zinc. (^) Add to a test tube half full of the hydrochloric acid a few drops of nitric acid and of silver nitrate solution. The white, curdy precipitate 49 Experiments. is silver chloride. Filter part of the contents of the test tube, and ex- pose the precipitate to the sunlight. Describe the change which soon occurs. To the remaining contents of the test tube add ammonium hydroxide, and shake. Describe the result. Experiment 44. Tests for Hydrochloric Acid or a Chloride. (a) What is a simple test for hydrochloric acid gas or for concen- trated hydrochloric acid ? {b} What is the usual test for hydrochloric acid ? (c) Dissolve a little sodium chloride in a test tube half full of water, and apply the test designated in (). (Suggestions. See Exps. 12 (b) and 43 ().) COMPOUNDS OF NITROGEN. Experiment 45. Preparation of Ammonia. Materials : 1 5 grams lime, 15 grams ammonium chloride, 3 bottles, 2 glass plates, pneu- matic trough filled as usual, litmus paper, stick of wood, filter paper. The apparatus is shown (in part) in Figure 106. The large test tube, A, is provided with a one-hole rubber stopper to which is fitted the right-angle bend, C, connected with a short glass tube, B (12 centimeters or 5 inches long), by the rubber tube, D. (a) Weigh the lime and ammonium chloride separately, mix them thoroughly on a piece of paper, and slip the mixture into the test tube to which a little water has been previously added. Add a little water. Quickly insert the stopper with its tubes, and clamp the test tube as shown in the figure (taking care not to crush the test tube) . Slip the glass delivery tube, B, into a bottle, invert the bottle, and hold it so that the tube is in the position shown in the figure. Heat the test tube gently with a low flame, beginning near the top of the mix- ture and gradually working downward. Ammonia gas will pass up into the bottle, which should be removed when full and covered with a glass plate. A piece of moist red litmus paper held near the mouth will show when the bottle is full. Do not smell at the mouth of the bottle. Collect two bottles and set aside until needed. FIG. 106. Appara- tus for preparing and collecting ammonia gas. Compounds of Nitrogen. 491 (b) As soon as the last bottle has been collected, rearrange the appa- ratus to absorb the ammonia gas in water, as in the case of hydrochloric acid (see Exp. 41 (^)). Replace the short glass tube by the delivery tube, E, which should pass through the wooden block, f y into a bottle, G, one fourth full of water, so that the end is just above the surface of the water (see Fig. 104). Continue to heat the generator at intervals, and the gas will be absorbed by the water. Shake the bottle occasionally. While the solution is being prepared, study the gas already collected. Experiment 46. Properties of Ammonia Gas. Proceed as follows with the ammonia gas prepared in Experiment 45 :- (a) Test the gas in one bottle with moist litmus paper and with a blazing stick. Describe the result. Compare the action with the behavior of hydrogen, oxygen, and hydrochloric acid gas, under similar circumstances. () Invert the same bottle and stand it upon the shelf of the pneu- matic trough. Describe any change noticed inside the bottle. What property of the gas is revealed ? Is it a marked property ? Test the contents of the bottle with litmus paper (both colors). (c) Pour a few drops of concentrated hydrochloric acid into an empty, warm, dry bottle. Roll the bottle until the inside is well coated. Cover it with a glass plate, invert it, and stand it upon a covered bottle of ammonia gas. Remove both plates at once, and hold the bottles together by grasping them firmly about their necks. Describe the action, giving all the evidence of chemical action. What is the white product ? Experiment 47. Properties of Ammonium Hydroxide. Remove the bottle in which the ammonia gas is being absorbed (see Exp. 45, ()), and study the resulting ammonium hydroxide as follows : , (a) Determine the general properties, e.g. taste and odor (cau- tiously), feeling, action with litmus. (b) Warm a little in a test tube. What gas is evolved? (c) Try the effect of ammonium hydroxide on a grease spot. Describe the result. Experiment 48. Neutralization of Ammonia. Materials: Ammonium hydroxide, hydrochloric acid, evaporating dish, sodium hydroxide solution, litmus paper, gauze-covered ring. 492 Experiments. Fill an evaporating dish one fourth full of ammonium hydroxide, and slowly add dilute hydrochloric acid, stirring constantly, until the solution is just neutral or faintly acid. Evaporate to dry ness, very slowly, on a gauze-covered ring. Test the residue as follows : (#) Is it an acid, alkali, or salt ? (^) Warm a little with sodium hydroxide solution. What is formed? Draw a conclusion as to the nature of the residue. (c} Support the dish on the gauze and warm gently until a decided change occurs. Describe the result. What compound do the fumes suggest? (y convenient size, is connected as shown in the figure, and is to be one third full of lime- water. The tube, ZT, is to be connected with a delivery tube passing into a pneumatic trough arranged to collect a gas over water. Fill AA' two thirds full of coarsely powdered soft coal, which should be held in place with a loose plug of shredded asbestos. See that all connections are gas-tight by heating the ignition tube gently ; if the ipparatus is tight, the expanded air will bubble through the bottle D. Readjust, if necessary. 506 Experiments. Heat the whole ignition tube gently at first, and gradually increase the heat, but avoid heating either end very hot, otherwise the closed end may soften and burst or the rubber stopper may melt. As the heat increases, watch for marked changes in B, CC', and D. As soon as the slow bubbling shows that all air has been driven out of the apparatus, collect, as previously directed, two bottles of the gas evolved. Cover the bottles with wet filter paper as soon as they are removed from the trough. When the last bottle has been removed, disconnect the ap- paratus at any convenient point between A' and C. Let the ignition tube cool. Test the gas by holding a lighted match near the mouth of a bottle. Observe and record the color and heat of the flame. Is smoke formed ? Repeat with the remaining bottle, and observe more closely any facts suggested, but not clearly shown, by the first observations. Examine the contents of the ignition tube. Does it resemble coke or some form of carbon ? Examine the bottle, B, for tarry matter. Does the paper in C show the formation of ammonia ? If the paper in C is black or brown, it is caused by lead sulphide, which is formed by the interaction of hydrogen sulphide and a lead compound. Did the gas contain hydrogen sulphide ? Did the bottle, Z), show the formation of carbon dioxide ? EXERCISES FOR THE CLASS: (1) Describe briefly the whole experiment. (2) Sketch the apparatus. (3) Summarize the properties of coal gas. Experiment 74. Combustion of Illuminating Gas. Materials : Pointed glass tube (see Int. 3 (V)), bottle, limewater. Attach a pointed glass tube to the rubber tube connected with the gas jet, and lower a small flame into a cold, dry bottle. Observe at once the most definite result inside the bottle. Remove and extinguish the flame, add a little limewater to the bottle, and shake. What are the two products of the combustion of coal gas ? What do the observa- tions show about the composition of the main constituents of coal gas ? Experiment 75. Construction of a Bunsen Burner. Take apart a Bunsen burner and study the construction. Write a short description of the burner. Sketch the essential parts. Carbon. 507 Experiment 76. Bunsen Burner Flame. Materials: Glass tube, powdered wood charcoal, pin, copper wire, wire gauze. I. (a) Close the holes at the bottom of a Bunsen burner and hold a glass tube in the upper part of the flame. Note the black deposit. What is it? Where did it come from? Open the holes and hold the 'blackened tube in the colorless flame. What becomes of the deposit? How is the flame changed, if at all? What does the experiment suggest about the luminosity of flame ? (b) Dip a glass tube a short distance into powdered wood charcoal, place the end containing the charcoal in one of the holes at the bottom of the burner, and blow gently two or three times into the other end. Describe and explain the result. Does it verify the answer to the last question in (a) ? (^) Open and close the holes of a lighted burner several times. Describe the result. Pinch the rubber tube to extinguish the flame, then light the gas at the holes. What change is produced in the flame? What causes the change? ANSWER : (1) What is the object of the holes? (2) Why does the gas burn at the top and not inside of the burner? (3) Why does the flame sometimes "strike back" and burn inside? (4) Why is the Bunsen flame nonluminous? II. (a) Hold a match across the top of the tube of a lighted Bunsen burner. When it begins to burn, Note where it is charred, and explain the result, a piece of wire down upon the remove and extinguish it. JL Press gauze flame. Describe the appearance of the gauze. The same fact may be shown by sticking a pin through a (sul- phur) match, suspending it across the burner, and then lighting the gas. The position of the match is shown in Figure log. Turn on a full FIG. no. Bent tube for ex- . , . amining the structure of a Bun- current of g as before sen flame. lighting it. What does the whole FIG. 109. Sul- phur match sus- pended across the top of a Bunsen burner. 508 Experiments. experiment show about the structure of the lower part of the Bunsen flame? Verify your answer by (b). (&) Bend a glass tube about 15 centimeters (6 inches) long into the shape shown in Figure no. Hold the shorter arm in the flame about 2 centimeters (i inch) from the top of the burner tube. Hold a lighted match for an instant at the upper end of the tube. What' does the result show about the structure of^he Bunsen flame? Does it verify (a) ? (c) Find the hottest part of the flame, when a full current of gas is burning, by holding a copper wire in the flame. Measure its distance, approximately, from the top of the burner tube. (d?) Examine a typical Bunsen flame one which shows clearly the outlines of the inner part. What is the general shape of each main part? Draw a vertical and a cross section of the flame. Experiment 77. Candle Flame. Materials : Candle, two blocks of wood, bottle, piece of stiff white paper, limewater, matches, lamp chimney, copper wire ( 15 centimeters or 6 inches long). Attach a candle to a block of wood by means of a little melted candle wax, and proceed as follows : (a) Hold a cold, dry bottle over the lighted candle. Describe the result produced inside the bottle. What is the product? What is its source? Remove the bottle, pour a little limewater into it, and shake. Describe and explain the result. What are the two main products of a burning candle? (b} Blow out the candle flame, and immediately hold a lighted match in the escaping smoke. Does the candle relight? Why? What is the general nature of this smoke? How is it related to the candle wax? How does (b} contribute to the explanation of (a) ? (c) Press a piece of stiff white paper for an instant down upon the candle flame almost to the wick. Repeat several times with different parts of the paper. What does the paper show about the structure of the flame? (d) Stand a lamp chimney over the lighted candle. How is the flame effected? Hold the chimney a short distance (i centimeter or .5 inch) above the block. Does the candle continue to burn? Why? Keep the chimney in the same position and cover the top with a block of wood. What is the result ? Why ? Carbon. 509 (V) Roll one end of the copper wire around a lead pencil to form a spiral about (2 centimeters or I inch) long. Press the spiral down upon the candle flame. What is the result? Why ? EXERCISES : (1) Draw a candle flame, showing the parts. (2) What is the essential difference between a candle flame and a Bunsen flame ? (3) Is there any essential difference between a candle flame and a gas or a lamp flame ? (4) Why do candles and lamps often smoke? Experiment 78. Kindling Temperature. () Press a wire gauze down upon a Bunsen flame. Where is the flame ? Hold a lighted match just above the gauze. Now where is the flame ? (<$) Extinguish the flame. Turn on the gas, hold the gauze in the escaping gas, about 5 centimeters (2 inches) above the top of the burner, and thrust a lighted match into the gas above the gauze. Where is the flame ? Lower the gauze slowly and describe the final result. (c} Hold the gauze in the flame in one position for a minute or two. Where is the flame at the end of this time ? Why ? EXERCISES : (1) Define kindling temperature. (2) What application is made of the principle illustrated by this experiment ? (3) State exactly how this experiment illustrates kindling tempera- ture. Experiment 79. Reduction and Oxidation with the Blow- pipe. Materials : Blowpipe, blowpipe tube, charcoal, lead oxide (litharge), sodium carbonate, sodium sulphate, wood charcoal, silver coin, zinc, lead, tin. Slip the blowpipe tube into the burner, light the gas and lower the flame until it is about 4 centimeters (1.5 inches) high. Rest the tip of the blowpipe on the top of the tube, placing the tip just within the flame. Put the other end of the blowpipe between the lips, puff out the cheeks, inhale through the nose, and exhale into the tube, using the cheeks some- what as a bellows. Do not blow in puffs, but produce a continuous flow of air by steady and easy inhaling and exhaling. The operation is nat- Experiments. ural and simple, and, if properly performed, will not make one out of breath. The flame should be an inner blue cone surrounded by an outer and almost invisible cone, though its shape varies with the method of production (see Fig. 44). Practice until the flame is produced volun- tarily and without exhaustion. Watch the flame and learn to distin- guish the two parts, so that they may be intelligently utilized. I. Reduction, (a) Make a shallow hole at one end of the flat side of a piece of charcoal. Fill the hole with a mixture of equal parts of pow- dered sodium carbonate and lead oxide, and heat the mixture in the reducing flame. The sodium carbonate melts and assists the fusion of the oxide, but the former is not changed chemically. In a short time bright, silvery globules will appear on the charcoal. Let the mass cool, and pick out the largest globules. Put one or two in a mortar, and strike with a pestle. Are they soft and malleable, or brittle and hard ? State the result when a globule is drawn across or rubbed upon a white paper. How do the properties compare with those of metallic lead ? What has become of the oxygen ? Of what chemical use is the charcoal ? (b) Grind together in a mortar a little sodium sulphate and wood charcoal, adding at intervals just enough water to hold the mass to- gether. Heat this paste fora few minutes in the reducing flame as in (a) . Scrape the fused mass into a test tube, boil in a little water, and put a drop of the solution on a bright silver coin. If a dark brown stain is produced, it is evidence of the formation of silver sulphide. Repeat, if no such stain is produced. State all the chemical changes which led to the production of the silver sulphide, explaining at the same time how the experiment illustrates reduction. II. Oxidation, (a) Heat a small piece of zinc on charcoal in the oxidizing flame. What is the product ? Observe its color, and the color of the coating on the charcoal when hot and cold. Record as described \&(d). (b) Heat a piece of lead as in (a}. Observe the presence or absence of fumes, as well as the color of the coating when hot and cold. See (d). (c) Heat a small piece of tin in the oxidizing flame. Observe as in (b} . (d} Tabulate the above observations, stating (i) the color of the hot and cold coating on the charcoal, (2) presence or absence of fumes, (3) name of product. EXERCISES : (1) Sketch a blowpipe. (2) Sketch a flame showing the oxidizing and reducing parts. Fluorine, Bromine, and Iodine. 511 FLUORINE, BROMINE, AND IODINE. Experiment 80. Preparation and Properties of Hydrofluoric Acid. Materials : Lead dish, glass plate, paraffin, file, calcium fluoride, concentrated sulphuric acid. Precaution. Hydrofluoric acid gas is a corrosive poison. An aque- ous solution of the gas commercial hydrofluoric acid burns the flesh frightfully. Warm a glass plate about 10 centimeters (4 inches) square by dipping it into hot water or by standing it near a warm object, such as a radiator. If it is held over a flame, it is liable to crack. Coat one surface with paraffin. The surface should be uniformly covered with a thin layer. Scratch letters, figures, or a diagram through the wax with a file. Be sure the instrument removes the wax through to the glass, and that the lines are not too fine. Put 5 grams of calcium fluoride in a lead dish and add just enough concentrated sulphuric acid to form a thin paste. Stir the mixture with a file. Place the glass plate, wax side down, upon the lead dish and stand the whole apparatus in the hood for several hours, or until some convenient time. Remove the plate. Scrape the contents of the dish, immediately, into a waste jar in the hood, and wash the dish free from acid. Most of the wax can be scraped from the glass plate with a knife. The last portions can be removed by rubbing with a cloth moistened with alcohol or turpentine. Do not attempt to melt off the wax over the flame. If the experiment has been properly performed, the plate will be etched where the glass was exposed to the hydrofluoric acid gas. Experiment 81. Preparation and Properties of Bromine. Materials : Potassium bromide, manganese dioxide, dilute sulphuric acid, bottle of water, test-tube holder. The apparatus is shown in Figure 1 1 1 . The large test tube is provided with a one-hole rubber stopper to which is fitted the bent glass tube. The latter is about 30 centimeters (12 inches) long, and is bent according to the directions given in the Introduc- paratus fo r pre . tion, 3 (b). paring bromine. 512 Experiments. Precaution. Bromine is a corrosive liquid which forms, at the ordinary temperature, a suffocating vapor. Perform in the hood all experiments which use or evolve bromine. - Put a dozen crystals of potassium bromide in the test tube, add an equal quantity of manganese dioxide and 10 cubic centimeters of dilute sulphuric acid. Insert the stopper and its tube securely, and boil gently. Do not hold the test tube in the hand, but use the test tube holder. Brown fumes soon appear in the test tube and pass out of the delivery tube. Regulate the heating so that this vapor will condense and collect in the lower bend of the delivery tube. Both vapor and liquid are bromine. When no further boiling produces bromine vapor in the test tube, pour the bromine from the delivery tube into a bottle of water. Observe and record the physical properties of this bromine, especially the color, solubility in water, specific gravity, volatility, and physical state. Try the action of the contents of the bottle on litmus paper ; if the action is not marked, push the paper down near the bro- mine. Determine the odor by smelling cautiously of the water in the bottle. As soon as these observations have been made, pour the con- tents of the bottle into the sink and flush with water, or pour into a jar in the hood. Wash the test tube free from all traces of bromine, taking care to get none on the hands. ANSWER : (1) In what ways does bromine physically resemble chlorine? In what ways does it differ from chlorine? (2) How is it essentially different from all other elements previously studied ? Experiment 82. Properties of Potassium Bromide. Materials : Potassium bromide, silver nitrate solution, ammonium hydroxide. Examine a crystal of potassium bromide, and state its most obvious properties. Dissolve it in a test tube half full of water, and add a few drops of silver nitrate solution. Describe the result. Is the solid prod- uct soluble in ammonium hydroxide? How can bromides be distin- guished from chlorides ? Do the properties of bromides, typified by potassium bromide, suggest any marked relation to chlorides ? Experiment 83. Preparation and Properties of Iodine. Ma- terials: Potassium iodide, manganese dioxide, mortar and pestle, con- centrated sulphuric acid, funnel, cotton. Fluorine, Bromine, and Iodine. 513 Grind together in a mortar a dozen large crystals of potassium iodide and about twice the bulk of manganese dioxide. Put the mixture in a test tube provided with a holder, moisten with water, and add a few cubic centimeters of concentrated sulphuric acid. Plug with cotton the inside opening of a funnel, and hold the latter firmly over the mouth of the test tube. Heat the test tube gently with a low flame (5 centime- ters or 2 inches). The vapor of iodine will fill the test tube, and crys- tals will collect in the upper part of the test tube and in the funnel. If the crystals collect in the test tube, a gentle heat will force them into the funnel. Continue to heat until enough iodine collects in the funnel for several experiments. Scrape the crystals into a dish. Study the properties as follows : () Observe and record the physical properties of iodine, especially the color of the solid and of the vapor, volatility, and odor (cautiously). () Heat a crystal in a dry test tube, and when the tube is half full of vapor, invert it. What does the result show about the density of iodine vapor? (c) Touch a crystal with the finger. What color is the stain ? Will water remove it? Will alcohol? Will a solution of potassium iodide? What do these results show about the solubility of iodine ? (NOTE. If crystals are left, use them in the next experiment. Pre- serve in a stoppered bottle.) Experiment 84. Test for Iodine with Carbon Bisulphide. Materials : Iodine, potassium iodide, carbon disulphide, chlorine water. Precaution. Carbon disulphide is inflammable. It should not be used near flames. (a) Free iodine. Add a few drops of carbon disulphide to a very dilute solution of iodine, made by dissolving a crystal of iodine in a solution of potassium iodide, and observe the color of the carbon disul- phide, which, being much heavier than water, will sink to the bottom of the test tube. How does it resemble the color of iodine vapor ? (b) Combined iodine. Add a few drops of carbon disulphide to a very dilute solution of potassium iodide. Is there positive evidence of iodine ? Now add several drops of chlorine water, and shake. How does this result compare with the final result in (a) ? The result is due to the fact that chlorine liberates iodine from its compounds, and the iodine, being free, exhibits the characteristic color. Experiments. Experiment 85. Test for Iodine with Starch. Materials: Starch, mortar and pestle, iodine solution, potassium iodide, chlorine water. Grind a lump of starch in a mortar with a little water to creamy con- sistency. Pour this into about 100 cubic centimeters of boiling water, and stir the hot liquid. Allow it to cool, or cool it by holding the vessel in a stream of cold water, and then pour off the clear liquid. Use this cold starch solution to test for iodine. (a) Free iodine. Add a few cubic centimeters of the starch solution to a test tube nearly full of water, and then add a few drops of iodine solution. The deep blue color is due to the presence of a compound which is always formed under these circumstances, but the composition of which is unknown. If the color is black, pour out half of the liquid and add more water, or pour some of the liquid into a dish of water. (b) Combined iodine. Add a few cubic centimeters of the starch solution to a very dilute solution of potassium iodide. Is the blue com- pound formed ? Add a few drops of chlorine water, and shake. Com- pare with the final result in Experiment 84 (b) . Experiment 86. Detection of Starch by Iodine. Materials: Dilute solution of iodine (in potassium iodide), mortar and pestle, potato, rice, bread. Test the potato, rice, and bread for starch by grinding a little of each with water in a mortar, and then adding a few drops of the extract to a very dilute solution of iodine. State the result in each case. Experiment 87. Properties of Potassium Iodide. Materials : Potassium iodide, silver nitrate solution, ammonium hydroxide. Proceed with the potassium iodide as in Experiment 82. How can iodides be distinguished from chlorides ? Do iodides, typified by potassium iodide, suggest any marked relation to bromides and chlorides ? SULPHUR AND ITS COMPOUNDS. Experiment 88. Properties of Sulphur. (a) Examine a lump of sulphur, and state briefly its most obvious physical properties. (6) Optional. Weigh a lump of roll sulphur to a decigram. Slip it carefully into a graduated cylinder previously filled with water to a Sulphur and its Compounds. 515 known point about half full and note the increase in the volume of water. This increase in volume is equal to the volume of the sulphur. Calculate the specific gravity of sulphur from the observed data. (NOTE. Specific gravity equals weight in air divided by weight of equal volume of water.) Experiment 89. Amorphous Sulphur. Materials: Sulphur, old test tube, evaporating dish. Put a few pieces of roll sulphur in an old test tube. Heat carefully until the sulphur boils, and then quickly pour the contents of the test tube into a dish of cold water. This is amorphous sulphur. Note its properties. Preserve, and examine it after twenty -four hours. Describe the change, if any. Define amorphous, and illustrate it by this experiment. Experiment 90. Crystallized Sulphur. J Materials : Sulphur (roll and flowers), Hessian crucible, carbon disulphide, evaporating dish. (a) Monodinic. Fill a small Hessian crucible nearly full of roll sulphur. Support the crucible in the ring of an iron stand, and heat until all the sulphur is melted. Let it cool, and as soon as crystals shoot out from the walls just below the surface, pour the remaining melted sulphur into a dish of cold water. When the crucible can be handled without discomfort, crack it open lengthwise. Observe and record the properties of the crystals, especially the shape, size, color, luster, brittleness, and any other characteristic property. Allow the best crystals to remain undisturbed for a day or two ; then examine again, and record any marked changes. (b} Orthorhombic. Put 3 grams of flowers of sulphur in a test tube and add about 5 cubic centimeters of carbon disulphide remember the precaution to be observed in using this liquid (see Exp. 84). Shake until all the sulphur is dissolved, then pour the clear solution into an evaporating dish to crystallize. It is advisable, though not absolutely necessary, to stand the dish in the hood or out of doors, where there is no flame and where the offensive vapor will be quickly removed. Watch the crystallization toward the end, and, if perfect crystals form, remove them with the forceps (see Fig. 49). Allow the i See Appendix, 3 (3), (5). Experiments. liquid to evaporate almost entirely, then remove and dry the crystals. Examine them as in (a) and record their properties. EXERCISES : (1) Tabulate the essential results in (a) and (). (2) Make an outline sketch of an orthorhombic crystal of sulphur. Experiment 91. Combining Power of Sulphur. Materials : Sulphur, deflagrating spoon, bottle, iron powder, hydrochloric acid. (a) Set fire to a little sulphur in a deflagrating spoon, and lower the spoon into a bottle. Cautiously waft the fumes toward the nose, and observe and describe the odor. The product is a mixture of two oxides of sulphur. What does their formation show about the combin- ing power of sulphur ? () Repeat Experiment 34. Results similar to that in (<) are obtained with copper and other metals. Draw a general conclusion regarding the power of sulphur to combine with metals. Experiment 92. Sulphur and Matches. (a) Examine a sulphur match. Do you detect any sulphur ? Where? (b) Light a sulphur match, and observe the entire action, as far as the sulphur is concerned. Describe it. (c) What is the function of the sulphur in a burning match ? Experiment 93. Preparation of Hydrogen Sulphide. Mate- rials: Ferrous sulphide, dilute hydrochloric acid, three bottles, three glass plates, stoppered bottle, litmus paper. Use the same apparatus as in Experiment 38. Precaution. Hydrogen sulphide is a poisonous gas and has an offensive odor. It should not be inhaled. Perform in the hood all ex- periments evolving hydrogen sulphide. (a) Construct and arrange an apparatus like that shown in Figure 104. Fill the test tube, A, one third full of coarsely powdered ferrous sulphide, insert the stopper tightly, pour enough hydrochloric acid through the safety tube to cover the contents of the test tube. Hydrogen sulphide gas is rapidly evolved. If the evolution of gas slackens or stops, warm gently or add more hydrochloric acid. Collect three bottles, removing each as soon as full and covering with a glass plate. Set aside until needed. Sulphur and its Compounds. 517 () As soon as the last bottle of gas has been removed and covered, put in its place a bottle one fourth full of water. Adjust its height (by wooden blocks or by lowering the generator) so that the end of the delivery tube reaches to the bottom of the bottle. Continue to pass the gas into the water, by heating the test tube if necessary. The gas will be absorbed by the water, forming hydrogen sulphide water. Preserve it in a stoppered bottle for Experiment 95. Proceed at once with next experiment. Experiment 94. Properties of Hydrogen Sulphide Gas. Study as follows the hydrogen sulphide gas prepared in Experi- ment 93 : (#) Waft a little of the gas cautiously toward the nose, and describe the odor. This is characteristic of hydrogen sulphide, and is a decisive test. Has the gas color? (b) Test the gas from the same bottle with both kinds of moist litmus paper. Is it acid, alkaline, or neutral ? (c) Bring a lighted match to the mouth of the same bottle. Observe the properties of the flame as in previous experiments. Observe cau- tiously the odor of the product of the burned gas ; to what compound is the odor due? What, then, is one component of hydrogen sulphide ? (d) Burn another bottle of hydrogen sulphide and hold a cold bottle over the burning gas. What additional experimental evidence does this result give regarding the composition of hydrogen sulphide ? (V) Repeat any of the above with the remaining bottle of gas. EXERCISES : (1) Summarize the properties of hydrogen sulphide gas. (2) State the experimental evidence of its composition. Experiment 95. Preparations and Properties of some Sul- phides. Materials: Hydrogen sulphide water prepared in Experiment 93, clean copper wire, clean sheet lead, bright silver coin, lead oxide (litharge) ; solutions of lead nitrate, arsenic trioxide (in hydrochloric acid), tartar emetic, zinc sulphate. (a) Shake the bottle of hydrogen sulphide water prepared in Experi- ment 93 (or a similar solution), and hold successively at the mouth or in the neck of the bottle (i) a clean copper wire, (2) a bright strip of lead, and (3) an untarnished silver coin. Describe the result in each case. These compounds are sulphides of the respective metals. 518 Experiments. () Put a little litharge the brownish yellow oxide of lead in a test tube, cover it with hydrogen sulphide water, and warm gently. The product is lead sulphide. Describe it. Explain the change. (c) Add hydrogen sulphide water to lead nitrate solution. The product is lead sulphide. Observe the color. (*/) Proceed as in (c) with the arsenic solution. Observe the color of the arsenic sulphide. (e) Proceed as in (c) with the tartar emetic solution. Tartar emetic is a compound of antimony. Observe the color of the antimony sulphide. (_/") Proceed as in (c) with the zinc sulphate solution. Observe the color of the zinc sulphide. Experiment 96. Preparation of Sulphur Dioxide. Materials : Sodium sulphite, concentrated sulphuric acid, litmus paper, three bottles, two glass plates, stick of wood, pink flower. The apparatus is constructed, arranged, and used as in Experiment 41, with one excep- tion. The safety tube must be replaced by a 'dropping tube made thus . Cut off the top of a thistle tube about 2.5 centimeters (i inch) below the juncture of the stem and cup, slip a short rubber tube (5 centi- meters, 2 inches, long ) over one end of the stem, attach a Mohr ? s pinch- cock to the rubber tube, and connect the tube with the cup. (a) Put about 10 grams of sodium sulphite in the large test tube, cover with water, and insert the stopper with its tubes. Adjust the ap- paratus as shown in Figure 104. Fill the cup with concentrated sulphuric acid, open the pinchcock a little, and let the acid flow drop by drop upon the sodium sulphide. Sulphur dioxide gas is evolved and passes into the bottle, which should be removed when full, as previously described. Moist blue litmus paper held at the mouth of the bottle will show when the latter is full. Collect two bottles of gas, cover each with a glass plate, and set aside until needed. (b) As soon as the second bottle of gas has been removed and covered, put in its place a bottle one fourth full of water. Adjust its height (if necessary) by wooden blocks, so that the end of the delivery tube is just above the surface. Continue to add the acid drop by drop, at intervals, and the gas will be absorbed by the water. Shake the bottle occasionally. Meanwhile study the gas already collected. Sulphur and its Compounds. 519 Experiment 97. Properties of Sulphur Dioxide Gas. Proceed as follows with the gas prepared in Experiment 96 (#) : (a) Observe and state the most obvious physical properties, e.g. color, odor (cautiously), density. (b) Hold a blazing stick in a bottle of the gas. Will the gas burn or support combustion ? What previously acquired facts would have enabled you to predict this result ? (c) Pour water into the same bottle of sulphur dioxide until half full, cover with the hand, and shake. What is the evidence of solution ? Is the resulting liquid acid, alkaline, or neutral ? (d) Moisten a pink flower with a few drops of water, hang it in the remaining bottle of sulphur dioxide, holding it in place by putting the stem between the glass and a cork. Observe and describe any change in the color of the flower. What is this operation called ? Experiment 98. Properties of Sulphurous Acid. Test as follows the solution of sulphurous acid prepared in Experi- ment 96 () : (a) Taste cautiously, and describe the result. (b) Apply the litmus test, and state the result. (c} Pour a few drops of concentrated sulphuric acid into the bottle. What gas is liberated ? Experiment 99. Action of Sulphuric Acid with Organic Matter. Materials : Concentrated sulphuric acid, sheet of white paper, sugar, starch, stick of wood. (a) Write some letters or figures with dilute sulphuric acid on a sheet of white paper, and move the paper back and forth over a low flame, taking care not to set fire to the paper. As the water evaporates the dilute acid becomes concentrated. Observe and describe-the result. Paper is largely a compound of carbon, hydrogen, and oxygen, and the hydrogen and oxygen are present in the proportion to form water. Explain the general chemical change in this experiment. (b} Fill a test tube one fourth full of sugar, add an equal bulk of water, stand the test tube in the rack, and add cautiously several drops of concentrated sulphuric acid. If there is no decided result, add more acid. What is the black product ? Compare the final result with that obtained in Experiment 59 (). Is the chemical action the same in 520 Experiments. each experiment ? Are the statements made in (a) about paper also true of sugar ? (c) Repeat (), using powdered starch instead of sugar. Describe the result. How does the result resemble that in (b} and in Experi- ment 59 (a) ? Predict the components of starch. In what simple way may the prediction be verified ? (//) Stand a stick of wood in a test tube one fourth full of concen- trated sulphuric acid. Allow it to remain in the acid for fifteen minutes, then remove the stick and wash off the acid. Describe the change in the stick. Does it resemble that in (#), (), and (c), and in Experiment 59? Experiment 100. Test for Sulphuric Acid and Sulphates. Materials: Sulphuric acid, sodium sulphate, barium chloride solution, calcium sulphate, charcoal, powdered charcoal, blowpipe, silver coin. (a) Repeat Experiment 12 (c) with sulphuric acid and with sodium sulphate solution. () Repeat Experiment 79 I () with calcium sulphate instead of sodium sulphate. EXERCISES : (1) State briefly the test for sulphuric acid and soluble sulphates. For insoluble sulphates. (2) How can a sulphate be distinguished from a sulphite ? SILICON AND BORON. Experiment 101. Preparation and Properties of Silicic Acid. Materials: Sodium silicate solution, hydrochloric acid, evaporating dish, gauze-covered ring. Add dilute hydrochloric acid to a test tube half full of sodium silicate solution, and shake. The jellylike precipitate is silicic acid. Rub some between the fingers and describe the result. Evaporate the precipitate to dryness in a porcelain dish which stands upon a gauze- covered ring in the hood. As the mass hardens, stir it with a glass rod. Toward the end, add more hydrochloric acid and evaporate to complete dryness. Then heat strongly for five minutes. The residue is silicon dioxide mixed with chlorides of sodium and potassium. Rub some between the fingers or across a glass plate. Is any grit detected ? State the chemical changes which occur in changing sodium silicate into silicon dioxide. Silicon and Boron. 521 Experiment 102. Tests with Borax Beads. Materials: Pow- dered borax, platinum test wire (see Int. 5 (4)), solutions of cobalt nitrate and copper sulphate, manganese dioxide. Make a small loop on the end of the platinum test wire, moisten it, and dip it into powdered borax. Heat it in the flame, rotating it slowly ; at first the borax swells, but finally shrinks to a small, transparent bead. If the bead is too small add more borax and heat again. After use, the bead may be removed by dipping it, white hot, into water ; the sudden cooling shatters the bead, which may then be easily rubbed or scraped from the wire. (a) Cobalt Compounds. Touch a transparent borax bead with a glass rod which has a drop of cobalt nitrate solution on the end. Heat the bead in the oxidizing flame. Observe the color when cold. If it is black melt a little more borax into the bead ; if faintly colored, moisten again with the cobalt solution. The color is readily detected by look- ing at the bead against a white object in a strong light, or by examining it with a lens. When the color has been definitely determined, heat again in the reducing flame. Compare the color of the cold bead with the previous observation. (b) Copper compozinds. Make another transparent bead, moisten it with copper sulphate solution and heat it first in the oxidizing flame, and then in the reducing flame. Compare the colors of the cold beads, and draw a conclusion. (c) Manganese Compounds. Make another transparent bead, touch it with a minute quantity of manganese dioxide, and proceed as in (b). Compare the colors of the cold beads, and draw a conclusion. (d) Tabulate the results of this experiment. EXERCISE : Draw a Bunsen flame, showing the reducing and oxidizing parts. Experiment 103. Preparation and Properties of Boric Acid and the Test for Boron. Materials: Borax, alcohol, evaporating dish, concentrated hydrochloric acid. To a test tube half full of boiling water, add about 10 grams of powdered borax. Add about 5 cubic centimeters of concentrated hydrochloric acid to this hot solution, and let the whole cool. Crystals of boric acid will separate. Filter. Describe the crystals. Put some of the crystals in an evaporating dish, add a little alcohol, 522 Experiments. and set fire to the solution. Observe the color of the flame. It is caused by a complex compound of boron, and is the test for this element. PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH. Experiment 104. Some Properties of Phosphorus. (a) Smell of the head of a phosphorus-tipped match. Describe the odor. (b) Rub the head of a phosphorus-tipped match in a dark place, and observe and describe the result. (c) The most striking property of phosphorus is the readiness with which it lights and burns in air. This property is too dangerous to try in the laboratory. Read about it in the text book. What application is made of this property ? Why ? Experiment 105. Test for Arsenic. Repeat Experiment 95 (, E, in Fig. 104) which reaches to the bottom of a small bottle; the latter has a two-hole stopper. The delivery tube passes 544 Experiments. through one hole, and through the other passes a bent tube connected with a U-tube. I. Put a liter of water in the bottle, add 150 grams of grape sugar, and shake until dissolved; pour 150 cubic centimeters of yeast into this solution. Fill the small bottle half full of limewater. Fill the U-tube with pieces of sodium hydroxide. Connect the apparatus and stand it in a dark place, where the temperature is 25-3o C. Fermentation begins at once, and carbon dioxide one of the prod- ucts bubbles through the limewater, which is protected from the action of the air by the sodium hydroxide. Examine the stopper for a leak, if no change occurs in the limewater. The operation should be allowed to continue at least a day, and longer if possible. The flask will then contain mainly water, unchanged grape sugar, alcohol, and some products of minor importance. Pour off the liquid, agitate it with a little bone black to remove the odor and color, and filter. The alcohol, which varies in quantity with the conditions, is dissolved in a large excess of water and must be separated by distillation. II. The distillation is performed with the apparatus used in Experi- ment 13. Fill the flask half full of the liquid from I, add a few pieces of pipestem (or granulated zinc, or glass tubing) to prevent " bumping," and distil about 50 cubic centimeters. Save the distillate. Replace the residue in the flask by nrore liquid from I, distil again, and repeat this operation until all the liquid has been used. Replace the one-hole stopper with a two-hole stopper, insert a thermometer in one hole so that the bulb just touches the surface of the combined distillates, which should now be distilled. Heat gently, and collect in a separate receiver the distillate which is formed when the liquid boils between 80 and 93 C. This distillate contains most of the alcohol. Test as follows : (a) Note the odor. ($) Drop a little into a warm dish, and hold a lighted match over it. If it does not burn, it shows that the alcohol is too dilute. Put a little in a dish, warm gently, and light the vapor. Describe the result. Experiment 166. Properties of Alcohol. (Optional. ) Materials : Alcohol, camphor, shellac, rosin, porcelain dish. () Determine cautiously the odor and taste of alcohol. Drop a little on a glass plate or on a piece of paper, and watch it evaporate. Is its rate of evaporation more rapid than that of water? Organic Compounds. 545 () Weigh a measured quantity (about 25 cubic centimeters) of 95 per cent alcohol and calculate its specific gravity. (c) Alcohol dissolves many organic substances. Try camphor, pow- dered shellac, or rosin. Describe the result. Verify the solvent power of alcohol by adding water to the solutions. Describe the result. (d) Burn a little alcohol in a dish and observe the nature of the flame. What are the products of combustion? Experiment 167. Preparation and Properties of Aldehydes. Materials: Concentrated hydrochloric acid, ethyl alcohol, potassium dichromate solution, methyl alcohol, copper wire, forceps. (#) Acetic Aldehyde. Add a little concentrated hydrochloric acid and several drops of ethyl alcohol to several cubic centimeters of potas- sium dichromate solution. Warm gently, and observe the peculiar- smelling gaseous product. It is aldehyde vapor, aldehyde itself being a colorless, extremely volatile liquid, which boils at 20.8 C. (b) Formic Aldehyde or Formaldehyde. Put a few cubic centimeters of methyl alcohol in a test tube and stand the test tube in a rack. Wind a piece of copper wire into a spiral around a glass rod or lead pencil. Slip the spiral from the rod, grasp one end into the forceps, and heat the wire red-hot in the flame. Then quickly drop it in the methyl alcohol. The pungent vapor which is suddenly produced is largely the vapor of formaldehyde. Experiment 168. Properties of Ether. Materials : Ether, evapo- rating dish, glass plate, wax. Precaution. Ether vapor is easily ignited, and should never be brought near a flame. (a) Pour a little ether into a dish or test tube and observe the odor and volatility. Taste cautiously. Pour a drop upon a glass plate or a block of wood. How does its rate of evaporation compare with that of alcohol ? Pour a little upon the hand and describe the result. () Add a bit of wax to a few cubic centimeters of ether, and shake. The result is typical ; draw a conclusion. Experiment 169. Properties of Acetic Acid. Treat acetic acid as follows : (a) Taste (cautiously), and describe. (b) Test with litmus paper, and describe the result. 546 Experiments. (V) Warm a little in a test tube, and smell (cautiously) . Describe the odor. Experiment 170. Properties of Vinegar. (a) Show, experimentally, that vinegar contains acetic acid. (^) Repeat Experiment 60, using vinegar instead of indigo solution. Experiment 171. Test for Acetic Acid and Acetates. Cautiously add a few drops of concentrated sulphuric acid to equal (and small) volumes of acetic acid and alcohol. Shake and warm gently. The pleasant, fruitlike odor is due to the vapor of ethyl ace- tate, a volatile liquid which is always formed under these circumstances. (NOTE. This experiment is also a test for alcohol.) Experiment 172. Preparation and Properties of Acetates. Materials for (a) : Sodium carbonate, acetic acid, concentrated sul- phuric acid, alcohol, porcelain (or agate) dish. For (b) : Litharge, acetic acid, porcelain dish. Prepare one or both of the following acetates : (a) Sodium acetate. Dissolve 20 grams of sodium carbonate in 10 cubic centimeters of water in a porcelain (or agate) dish, and slowly add 30 cubic centimeter? of commercial acetic acid, with constant stir- ring. If the solution is not acid, add a little more acetic acid. Filter the solution, if not clear. Evaporate to crystallization. When the crystals have formed, remove and dry them. Describe the crystals. Prove that they contain water of crystallization. Test the acetate as follows: (i) Dissolve a little in water, add a few drops of concen- trated sulphuric acid, and boil. What does the odor show is present? What other acids have been similarly prepared? (2) Dissolve as in (i), add a few drops of alcohol and of sulphuric acid, and boil. What does the odor conclusively prove? Preserve the crystals, finally, in a glass-stoppered bottle, or in one having a cork covered with paraffin. (b) Lead acetate {poisonous). To 10 grams of litharge add 18 cubic centimeters of commercial acetic acid in small portions. Stir the mix- ture constantly during the addition of acid. After all the acid has been added, heat gently until the action ceases. (If the solution is green or bluish, it is due to a copper compound. The copper may be precipi- tated and removed mechanically by standing a strip of lead in the solu- tion for an hour or more. Pour off the clearer liquid and then filter.) Organic Compounds. 547 Evaporate cautiously to crystallization. Remove the crystals from the liquid, and dry at a moderate temperature. Preserve the crystals finally as in (a). Describe the crystals. Test them for lead (see Exp. 145 ()), and for an acetate. Experiment 173. Properties of Certain Organic Acids. Mate- rials: Tartaric and citric acids, potassium permanganate solution, sodium bicarbonate, sugar, concentrated nitric acid, evaporating dish, litmus paper. (1) Tartaric acid. Observe and describe the results in the follow- ing : (a) Taste cautiously a dilute solution of tartaric acid, (b) Apply the litmus test, (c) Add a little of the solution to a sodium bicarbonate solution. (W) Dissolve two or three crystals of potassium permanga- nate in a test tube half full of water, add a little sodium hydroxide solu- tion and two or three pieces of tartaric acid (solid). Warm gently, but do not shake. The change is due to the reduction of the potassium permanganate by the tartaric acid. (2) Citric acid. Proceed as in (i) with citric acid. (3) Oxalic Acid, (a) This acid is poisonous. Do not taste it. (ti) and (c) Proceed as in (i). (*/) Dissolve two or three crystals of potassium permanganate in a test tube half full of water and add half the volume of sulphuric acid. Add oxalic acid solution until a decided change appears. Describe and explain it. (e) Add a few drops of ink to oxalic acid solution, and shake. Describe the result. Experiment 174. Preparation and Properties of Ethyl Ace- tate. Repeat Experiment 171. ANSWER : (1) What class of organic compounds does ethyl acetate represent ? What general property has this class ? (2) To what inorganic compound does ethyl acetate correspond? (3) What is the relation of ethyl acetate to (a) alcohol and (^) ace- tic acid ? Experiment 175. Preparation of Soap. Prepare soap in an iron or a tin dish by one of the following methods : 548 Experiments. (a) Dissolve 10 grams of sodium hydroxide in 75 cubic centimeters of water, add 30 grams of lard, and boil until the mixture begins to solidify. Then add 20 grams of fine salt in small portions. Stir constantly during the addition of the salt. Boil a few minutes. Let the mass cool, and then remove the soap, which will form in a cake at the surface. (#) Dissolve 13 to 15 grams of sodium hydroxide in 100 cubic centi- meters of water, add 100 cubic centimeters of castor oil, and boil for about half an hour. Add 20 grams of salt, and then proceed as in (#). (c) Dissolve 8 grams of potassium hydroxide in 150 cubic centi- meters of alcohol, add 10 grams of lard, and stir constantly while the mixture is being heated cautiously to sirupy consistency. Allow the solution to cool. The jellylike product is soap. Preserve a sample. Experiment 176. Properties of Soap. Materials: Soap, sul- phuric acid, calcium sulphate, magnesium sulphate, and acid calcium carbonate solutions. Test as follows the soap prepared in Experiment 175 : (#) Leave soap shavings exposed to the air for several days. What does the result show about the presence of water in the soap ? () Test soap solution with litmus paper. (c) Add considerable dilute sulphuric acid to a soap solution. The precipitate is a mixture mainly of palmitic and stearic acids. Describe it. (d) To a little soap solution in separate test tubes add calcium sul- phate and magnesium sulphate solutions. Describe the result. Boil for a few minutes and describe the result. Prepare a solution of acid calcium carbonate by passing carbon dioxide into limewater until the precipitate is redissolved (see Exp. 69). Add some of the solution to a soap solution, and describe the result. Boil, as above, and describe the result. ANSWER : (1) -What is hard water ? Soft water ? (2) What is permanent hardness? Temporary hardness? How can the later be removed ? Experiment 177. Properties of Glycerine. O) Add a little glycerine to a test tube half full of water, and shake. Add considerable more glycerine, and shake. What does the result show about the solubility of glycerine in water? Laboratory Equipment. 549 () Cautiously taste the liquid resulting from (a). Describe the result. Experiment 178. Fehling's Test for Sugar. Materials: Copper sulphate, Rochelle salt, sodium hydroxide, and grape sugar solutions. Mix equal (and small) volumes of copper sulphate, Rochelle salt, and sodium hydroxide solutions in a test tube, and boil carefully. The mixture should be strongly alkaline. Add a little grape sugar solution, and boil until a decided change is produced. The precipitate is cuprous oxide. Describe it. (NOTE. Cane sugar must be changed to grape sugar by boiling with dilute sulphuric acid before the above test is applicable.) Exercises for Review. 1. What happens to sugar and starch (a) when heated, and () when treated with concentrated sulphuric acid ? 2. What is the test for starch ? 3. Discuss the solubility of alcohol in water. 4. What is the effect of heat upon paper and cotton? Of potassium permanganate on paper? Experiment 179. Properties of Benzene. Put one or two drops of benzene in an evaporating dish, and cautiously bring a lighted match near it. Describe the result. LABORATORY EQUIPMENT. The Equipment of a laboratory should be limited solely by the means at the disposal of the teacher. Accurate and rapid work is largely determined by the available facilities, and no pains should be spared to secure the equipment which will yield the largest educational' return for the time and money expended. The lists given below include the apparatus and chemicals needed for the experiments in this book. Quantities and prices have been omitted in justice to teachers, dealers, and the author. Different teachers use different quantities, prices fluctuate, and qualities vary. The author, at his own suggestion, has lodged with the L. E. Knott Apparatus Co., 16 Ashburton Place, Boston, Mass., information regarding the quantities 550 Experiments. of apparatus and chemicals used by his classes. It is hoped that teach- ers will correspond with both author and dealer when preparing order lists. The author takes this opportunity to say that he has no financial connection whatever with any dealer in scientific supplies. LIST A. INDIVIDUAL APPARATUS. This list includes the apparatus constantly used by a single student, who should be provided with each piece. The set will cost from $4.75 to $5. The discount on apparatus in this and succeeding lists depends upon the total amount of the order. 6 Test tubes, 6 x |. 3 Test tubes, 8 x i . Test-tube holder. Test-tube rack. Test-tube brush. Bunsen burner. Blowpipe. Blowpipe tube. Bottles, wide mouth, 250 cc. Funnel, 2.\ in. Evaporating dish. Pair iron forceps. Triangular file. i Mortar and pestle, 3 in. i Deflagrating spoon. 1 Pneumatic trough. 2 ft. Rubber tubing, \ in. in diam. 100 Filter papers, 4 in. i ft. Glass rod. 6 in. Rubber tubing, T V i One-hole and i two-hole rubber stopper to fit large test tube. 4 ft. Glass tubing to fit rubber stoppers (above), i Safety tube. LIST B. SPECIAL APPARATUS. This list includes apparatus used occasionally. Numbers in paren- theses refer to experiments. The set will cost from $3 to $3.25. i Crucible, Hessian, 4 in. deep (59, 90). Dish, lead (80). Flask, Erlenmeyer, 250 cc. (25). Pinchcock, Mohr (96). Screw, Hofmann (25). Thistle tube (96) . Lamp chimney (77). i Graduated cylinder, 25 cc and others), i Magnet (156). i Candle (63, 77). i Sand-bath pan, 4 in. i Wing-top burner (Int. i Dish, iron or tin (109, in) i Retort, 250 cc. (49). (25 List E. Chemicals. 55 1 LIST C. APPARATUS FOR TEACHER'S EXPERIMENTS. This list includes the additional apparatus for the Teacher's Experi- ments. Numbers in parentheses refer to experiments. The set will cost about $11. Electrolysis apparatus (22, 36). Flask, 500 c.c. (10, 13, 165). Two-hole rubber stopper for above. U-tube (10, 73, 165). 2 One-hole rubber stoppers for above. 4 in. Capillary tubing (10, 22, 36). 3 ft. Glass tubing to fit rubber stoppers. i Safety tube (10). i Condenser complete (13, 165). i Tripod (13). i Thermometer (165). i Chlorine tube (23). i Ignition tube, 6 in. (73). i Bottle, wide mouth, 50 cc. (10). Battery, 3 cells (Grenet) (22, 36, 37). i Bottle, 2000 cc. (165). LIST D. GENERAL APPARATUS. This list includes the general laboratory apparatus. It should be extended as demands arise. It does not include such items as dupli- cate stoppers, extra glassware, tools, etc. Special inexpensive articles are noted in the experiments and in the "Handbook for Teachers" accompanying this book. Corks, assorted. Copper wire, No. 24. Glass plates, 4 x 4 in. Iron stands, 3 rings, 2 clamps. Matches. Wire gauze, iron, 4 x 4 in. Wooden blocks, 6 x 6 x i in., 6 x 6 x f in., 4 x 4 x in. (with I- in. hole in center see Exp. 38). Sand. Wood, thin sticks (Exp. 6 and others) . Rule, foot and 30 cm. Scales, trip. Weights for above. Tapers. Emery paper. Kerosene lamp. Graduated cylinders, 500 cc., 100 cc. LIST E. CHEMICALS. This list includes the chemicals needed for this book. Numbers in parentheses refer to experiments in which the chemicals are used. Acid, acetic, citric. Acid, hydrochloric, nitric. 55 2 Experiments. Acid, oxalic. pyrogallic (25). sulphuric, tartaric. Alcohol, ethyl. methyl (167). Alum, chrome. potassium. Aluminium, metal. sulphate. Ammonium, chloride, hydroxide, nitrate, oxalate. sulphide. Arsenious oxide (95, 105). Asbestos, shredded (73). Baking powder (68) . Barium chloride. nitrate (125). Benzene (179). Bismuth (107). Bleaching powder. Borax (powd.). Cadmium chloride (131). Calcium carbide (72). carbonate (marble), chloride, fluoride (80). oxide (lime), sulphate. Carbon disulphide. Chalk (native) (68). Charcoal, animal (powd.). lump. wood (powd.). Coal, soft. Cobalt nitrate. Cochineal. Coin (silver). Copper nitrate, sheet. sulphate (cryst.). Cotton (absorbent). Cream of tartar. Ether. Galena (148). Gelatine. Glycerine. Gold leaf (book). Hematite (161). Indigo. Iodine. Iron, chloride (#:). filings. powder. pyrites (161). sulphate (ous). sulphide (ous). wire (fine). wrought. Kerosene. Lead acetate. carbonate. dioxide (peroxide). nitrate. monoxide (litharge). sheet. tea. tetroxide. Limonite (161). Litmus paper. Magnesium oxide, ribbon, sulphate. Magnetite (161). Manganese dioxide, sulphate. Solutions. 553 Mercury. Mercuric chloride, nitrate, oxide (7). Mercurous nitrate. Mustard. Nickel chloride (162). Paraffin. Phenolphthalein (66). Picture cord (iron) (6). Potassium, metal (no). bromide. carbonate. chlorate (cryst.). chlorate (powd.). chloride. chromate. dichromate. ferricyanide. ferrocyanide. hydroxide. iodide. nitrate. permanganate. sulphate. sulphocyanide (thiocya- nate) (157-158). Pyrite (161). Rochelle salt (178). Rosin. Selenite (gypsum, cryst.). Shellac. Siderite (161). Silver nitrate. Soap. Soda lime (164). Sodium, metal. bicarbonate. carbonate. chloride. hydroxide. hyposulphite (thiosul- phate). nitrate. phosphate (disodium phosphate) . silicate (101). sulphate. sulphite (96). Stannous chloride (tin crystals). Starch. Steel. Strontium nitrate (123). Sugar, cane. grape (165). Sulphur, flowers. roll. Tartar emetic. Tin, granulated. Tooth powder (68). Turpentine. Vaseline. Vinegar. Water, distilled. Whiting (68). Wood ashes. Zinc, granulated, oxide, sheet, sulphate. SOLUTIONS. The following solutions are needed .for the experiments in this book. Those not included are described in the experiments requiring their use. 554 Experiments. Alum, 10 per cent. Ammonium chloride, 10 per cent. Ammonium hydroxide, i vol. to 3 vols. water. Ammonium oxalate, 1 4 per cent. Ammonium sulphide, i vol. to i vol. water. Barium chloride, 2 5 per cent. Battery solution (Grenet). Dis- solve 103 gm. powdered potas- sium dichromate in i liter of water and slowly add 103 gm. cone, sulphuric acid with con- stant stirring. Calcium chloride, 10 per cent. Chlorine water, 1 saturated (see Exp. 23, 38). Cobalt nitrate, 5 per cent. Cochineal. Prepare as described under Indigo. Copper sulphate, 10 per cent. Disodium phosphate, lo.per cent. Ferric chloride, 5 per cent. Ferrous sulphate, 1 10 per cent. Hydrochloric acid, i vol. to 4 vols. Indigo. Grind a little with water and dilute as desired. Iodine. Grind to solution 12 gm. iodine, 20 gm. potassium iodide, 10 cc. water, and add to 1000 cc. water. Lead acetate, 10 per cent. Lead nitrate, 10 per cent. Limewater. Let water stand over lime for several days, and siphon off the clear liquid. Magnesium sulphate, 10 per cent. Manganese chloride, 10 per cent. Mercuric chloride, 5 per cent. Poi- son. Mercurous nitrate, 8 5 per cent. Nitric acid, i vol. to 4 vols. water. Potassium bromide, 5 per cent. Potassium chloride, 5 per cent. Potassium chromate, 10 per cent. Potassium dichromate (or bichro- mate), 5 per cent. Potassium ferricyanide, 10 per cent. Potassium ferrocyanide, 10 per cent. Potassium hydroxide, 10 per cent. Potassium iodide, 5 per cent. Potassium nitrate, 10 per cent. Potassium permanganate, 2 5 per cent. Potassium sulphate, 10 per cent. Potassium thiocyanate (or sulpho- cyanide), i per cent. Silver nitrate, 5 per cent. Sodium carbonate, 10 per cent. Sodium chloride, 10 per cent. Sodium hydroxide, 10 per cent. Stannous chloride. 1 Dissolve 500 gm. of the salt in 1000 cc. hot cone, hydrochloric acid, and add a piece of tin. Sulphuric acid, i vol. to 4 vols. water. Tartar emetic, 10 per cent. Zinc sulphate, 10 per cent. 1 Must be freshly prepared. 2 Use distilled water. 8 Use distilled water, and add 75 cc. concentrated nitric acid and a little mercury. INDEX. Absolute zero, 439. Acetates, 417. Ethyl, 419. Metallic, 419. Acetic acid, 415. Constitution, 170, 416. Glacial, 415. Preparation, 415. Properties, 415. Series, 414. Test, 419. Acetylene, 116, 205. As illuminant, 206. Burner, 207. Composition, 206. Explosive properties, 205. Flame, 207, 216. Generation, 207. Liquid, 205. Series, 202. Acetylides, 206. Acid, acetic, 415. Benzoic, 431. Boracic, 261. Boric, 261. Butyric, 417. Capric, 421. Caproic, 421. Carbolic, 431. Carbonic, 194. Chloric, 91. Chlorous, 91. Citric, 419. Cyanic, 198. Ethyl sulphuric, 414. Fuming sulphuric, 251. Acid, continued. Gallic, 432. Glacial acetic, 415. Glacial phosphoric, 268. Hydriodic, 232. Hydrobromic, 230. Hydrochloric, 140. Hydrocyanic, 198. Hydrofluoric, 227. Lactic, 418. Malic, 418. Metaphosphoric, 268. Metastannic, 355. Muriatic, 92, 140. Nitric, 154. Nitrosylsulphuric, 248. Nitrous, 159. Nordhausen sulphuric, 252. Orthophosphoric, 268. Oxalic, 417. Palmitic, 417. Perchloric, 91. Picric, 431. Prussic, 198. Pyrogallic, 431. Pyroligneous, 415. Pyrophosphoric, 269. Pyrosulphuric, 252. Salicylic, 431. Silicic, 257. Stearic, 417. Sulphocyanic, 198. Sulphuric, 246. Sulphurous, 244, 245. Tannic, 432. Tartaric, 418. 555 556 Index. Acid calcium sulphate, 245. Of air, 196. Oxide, 97. Phosphate, 90, 269. Potassium fluoride, 226. Reaction, 90. Salt, 96. Sodium carbonate, 289. Sodium sulphate, 245. Sulphates, 251. Acidity, 94. Acids, 90. And ionization, 129. And oxygen, 18. Chlorine, 91. Commercial names, 92. Defined, 90. Dibasic, 92. General properties, 87. In familiar substances, 90. Monobasic, 92. Nomenclature, 91. Organic, 92, 414. Oxygen in, 97. Relation of oxides to, 96. Tribasic, 92. Addition products, 204. Agate, 255. Air, 61. 4f Acid of, 196. Alkaline, 149. Bad, 67. Composition, 64. Dephlogisticated, 18. Empyreal, 18. Fixed, 196. Gravimetric composition, 66. Hydrogen dioxide in, 60. Liquid, 69. Marine acid, 140. Mixture, 69. Relative humidity, 66. See Atmosphere. Slaked lime, 324. Solubility, 69. Air, continued. Volumetric composition, 64. Weight of liter, 62. Alabaster, 326. Alchemists, 154, 157, 158, 235, 246 251, 308, 314, 354, 3 7- Alcohol, ethyl, 409. Absolute, 410. Commercial, 410, 411. Constitution, 407. Fermentation, 416. Formulas, 407. Oxidation, 412. Preparation, 410. Pure, 410. Test, 419. Uses, 410. Alcohol, methyl, 409. Triacid, 420. Wood, 409. Alcoholic liquors, 411. Alcohols, 408. Constitution, 409. General nature, 408. Aldehyde, acetic, 412. Benzoic, 431. Formic, 412. Salicylic, 432. Aldehydes, 412. Alizirin, 367, 432. Alkali, 92, 93, 94. Action on litmus, 92. And glass, 259. Metals, 284. Sodium carbonate, 289. Volatile, 93, 149. Alkalies, common names, 92. Fixed, 93. Properties, 93. Alkaline, 92. Air, 149. Earth metals, 319. Reaction, 92. Silicate, 257, 258. Alkaloids, 433. Index. 557 Allotrope, 191. Allotropic modification, 191. Carbon, 190. Silicon, 255. Sulphur, 239. Allotropism, 190. Allotropy, 191. Alloys, 282. Antimony, 356, 360. Copper, 305. Fusible, 337, 360. Lead, 360. Magnesium, 346. Manganese, 369. Mercury, 339. Nickel, 306. Platinum, 394. Silver, 311. Tin, 356. Zinc, 306, 336. Allylene, 202. dels, 338. rn, 349. nium, 349. 349. Cake, 348. Chrome, 350, 367, 368. Concentrated, 349. General formula, 350. History, 350. Iron, 386. Potassium, 349. Shale, 350. Slate, 350. Sodium, 349. Alumen, 343. Alumina, 346. Preparation, 347. See Aluminium oxide. Aluminates, 348. Aluminium, 343. Acetate, 350, 417. Alloys, 344, 346. Bronze, 305, 346. Carbide, 203. Aluminium, continued. Chloride, 351. History, 343, 344, 345. Hydroxide, 348. Impurities, 345. In gems,- 347. Leaf, 346. Metallurgy, 343, 344. Name, 343. Occurrence, 343. Older processes, 344. Oxide, 343, 346, 347. Price, 345. Production, 344. Properties, 345. Silicate, 351. Sulphate, 348. Test, 347. Uses, 346. Alumino ferric cake, 348. Aluminum. See Aluminium. Alumium, 343. Alunite, 350. Amalgamated zinc, 339. Amalgamation, 281. Process for silver, 309. Amalgams, defined, 282, 339. 4fctd 339- Tin, 339, 356. Amethyst, 255. Oriental, 347. Ammonia, 147. Anhydrous, 148. As a refrigerant, 152. Composition, 153. Formation, 147. From coal, 147, 148. In ice-making, 152. Liquefied, 148, 149, 153. Muriate of, 151. Near stables, 147. Of commerce, 148. Preparation, 147. Properties, 148. Soda process, 289. 558 Index. Ammonia, continued. Uses, 152. Water, 149. Ammoniacal liquor, 213. Ammonium, 150. Alum, 349. Carbonate, 152. Chloride, 151. Chloroplatinate, 394. Compounds, 152. Hydroxide, 148, 149, 150. Molybdate, 369. Nitrate, 152. Sulphate, 151. Sulphide, 152. Sulphocyanate, 152. Salts, 150. Amorphous, 184. Carbon, 181, 184, 190. Sulphur, 239, 240. Amygdalin, 432. Amyl acetate, 419. Valerate, 419. Anaesthetic, 412, 413. Analysis, 3, 50. Qualitative, 242. Spectrum, 403. Water, 39. Anglesite, 357. Anhydride, 97. Carbonic, 195. Nitric, 163. Nitrous, 163. Anhydrite, 326. Anhydrous, 46. Aniline, 431. Dyes, 431. Animal charcoal, 189. Anion, 120, 121. Annealing glass, 260. Anode, 120, 121, 285, 291, 303, 312, 332, 344- Anthracene, 432. Anthracite coal, 185, 186. Antichlor, 138, 245, 252. Antidote for arsenic poisoning, 273, 385. Antifriction metals, 336. Antimony, 274. Acids, -274. Alloys, 356, 360. As metalloid, 278. Chlorides, 275. Compounds, 419. Name, 274. Oxides, 274. Oxychlorides, 275. Test, 275. Trisulphide, 274. Apatite, 225, 265. Aqua ammonia, 148, 149. Fortis, 158. Regia, 160, 316. Argentiferous lead, 308. Argentite, 308. Argentum, 308. Argol, 418. Argon, 68, 404. Aristotle, 61. Armor plate, 380, 389. Arrhenius, 128, 442. Arsenic, 272. Acids and salts, 273. Antidote, 273, 385. As metalloid, 278. Marsh's test, 273. Ores, 272. Oxide, 272. Poisoning, 273, 385. Production, 272. Pyrites, 272. Sulphide, 273. Test, 273. Trioxide, 272. Uses, 272. Vapor density, 169. White, 272. Arsenious oxide, 272. Arsenolite, 272. Arsine, 273. Index. 559 Artificial diamonds, 346. Graphite, 118. Stone, 333. Asbestos, 331. Ash, black, 288. Seaweed, 230, 231. Ashes and potassium compounds, 298. Atmosphere, 61. An, 62. And plants, 194. Argon in, 68. Carbon dioxide in, 67. Composition, 64. Inert gases in, 69. Ingredients, 62. Nitrogen in, 72. Of sun, 23. Oxygen and nitrogen in, 63. Ozone in, 22. Pressure, 62. Properties, 61. Water vapor in, 31, 66. See Air. Atomic theory, 79. Atomic weights, 101. And symbols, 103. And valence, 178. Classification by, 397. Determination, 170. Methods of determining, 173. Relation of properties to, 398. Standards, 102. Table, 448. Atoms, 79, 81. And ions, 125. And molecules, 80. Combining power, 176, 177, 406. In a molecule, 168, 174, 191, 204, 232, 238, 267, 272, 286, 336, 339, 407, 430. Replacement of, 176. Space relations, 178. Attraction, chemical, 4. Aurum, 314. Avogadro, 167, 442. Hypothesis, 167. Azote, 63. Azurite, 301, 308. Babbitt's metal, 336. Bacteria, 155. Baking powder, 290, 418. Soda, 290. Balard, 230, 442. Bamboo, 257. Barite, 329. Barium, 328. Carbonate, 329. Chloride, 329. Compounds, 329. Dioxide, 59. Nitrate, 329. Oxides, 12, 59, 329. Sulphate, 329, 362. Sulphide, 329. Test, 329. Barley sugar, 424. Baryta water, 329. Base, 92, 93. Ammonium hydroxide as, 150. Diacid, 94. Monacid, 94. Triacid, 94. Bases, 88. And ionization, 129. Nomenclature, 93. Relation of oxides to, 96. Basic, 93. Bismuth nitrate, 276. Oxides, 97. Salt, 96. Basicity, 92. Basil Valentine, 246. Bath metal, 305. Battery, electric, 120. Leclanche, 151. Baux, deposits at, 348. Bauxite, 348. Becher, 16, 442. 560 Index. Beef fat, 420. Beehive oven, 189. Beer, 193, 411. Beet sugar, 424. Potassium carbonate from, 297. Bell metal, 306. Bench of retorts, 210. Benzene, 213, 430. Constitution, 430. Derivatives, 430. Series, 202. Source, 429. Benzine, 208, 430. Benzoic acid, 431. Aldehyde, 431. Benzol, 430. Bergman, 196, 443. Berlin blue, 388. Berthollet, 77, 443. Beryllium, 401. Berzelius, 443. And Dulong, 57. Bessemer, 443. Bessemer steel, 381. Beverages, sparkling, 193. . Bicarbonate, sodium, 195. Binary compounds, 95, 176. Bismite, 275. Bismuth, 275. Carbonate, 275. Dioxide, 276. Hydroxide, 276. Nitrate, 276. Oxychloride, 276. Pentoxide, 275. Subnitrate, 276. Sulphide, 275. Test, 276. Trichloride* 276. Trioxide, 275. Bismuthinite, 275. Bismutite, 275. Bisulphite of soda, 245. Bittern, 229. Bituminous coal, 185, 186, 210, Bivalent elements, 176. Black, 196, 334, 443. Black ash process, 288. Damp, 203. Lead, 183, 357, 359. Magnesia, 370. Oxide of manganese, 370. Blast furnace, 275, 281. Lamp, 29. Bleaching by chlorine, 136, 137, 138. Hydrogen dioxide, 60. Sodium peroxide, 293. Sulphur dioxide, 244. Bleaching powder, 137. Block tin, 355. Pipe, 40. Blood and oxygen, 16, 17. Iron, 373. Blow, water gas, 213. Blowpipe, 222. Flame, 29, 222. Oxyhydrogen, 17, 28. Blue paint, 417. Print paper, 388. Stone, 307. Vitriol, 307. Bonds, 407. Bone ash, phosphorus from, 265. Cupel, 310. Bone black, 189. Bones, 271. Phosphorus from, 265. Books, reference, 450. Boracic acid, 261. Boracite, 261. Borax, 261, 262. And soldering, 263. Bead, 262. Boric acid, 261. Borides, 261. Bornite, 301, 373. Boron, 261. Bort, 182. Boyle, 64, 443. Law, 19. Index. S 6i Brand, 265. Brandy, 411. Brass, 305. Cyprian, 301. Braunite, 369. Bread making, 419, 427. Breathing, 16, 17. Bricks, 352. Brimstone, 238. Brines, 287. Britannia metal, 306, 356. British coal fields, 187. Gum, 427. Brittle metals, 279. Bromides, 230. Bromine, 228. Commercial process, 229. Compounds, 230. Discovery, 230. Name, 229. Production, 230. Properties, 229. Uses, 230. Water, 229. Bronze, 305. Aluminium, 305. Phosphor, 305. Silicon, 305. Brown iron ore, 374. Bullets, 360. Bunsen, 64, 219, 284, 403, 443. Burner, 219. Flame, 219, 220, 221. Burette, 98. Burner, acetylene, 207. Bunsen, 219. Self-lighting, 26. Burning, 15. Burnt alum, 349. Butane, 203. Butter, 421. Artificial, 421. Rancid, 417. Butylene, 204. Butyric acid, 417. Cadmium, 337. Sulphide, 337. Test, 337. Vapor density, 169. Caesium, 284, 299, 403. Caffeine, 433. Cake, alum, 348. Alumino ferric, 348. Press, 296. Salt, 288. Calamine, 334. Calcarone, 236. Calcination of ores, 281. Calcite, 320. Calcium, 319. Preparation, 319. Properties, 319, 320. Test, 328. Calcium and carbonate, 195. Acid sulphate, 429. Borate, 262. Carbide, 116, 205, 207. Carbonate, 192, 195, 319, 320, 321. Chloride, 67, 327, 328. Fluoride, 225, 227. Hydroxide, 325. See Limewater. Hypochlorite, 137. Iodide, 319. Light, 29. Magnesium carbonate, 331. Manganite, 370. Nitrate, 155. Oxide, 324. See Lime. Sulphate, 326, 327. Sulphide, 288, 328, 329. Calculations, chemical, 103. Calico printing, 350. Calomel, 340. Caloric, 112. Candle flame, 216, 217, 218. Power, 216. Candles, stearin, 422. Illuminating gas, 216. Cane sugar, 423. See Sugar. 562 Index. Cannizzaro, 443. Capric acid, 421. Caproic acid, 421. Caramel, 424. Carat, diamond, 183. Gold, 316. Carbide, aluminium, 203. Calcium, 116, 205, 207. Iron, 209. Carbohydrate, 423. Carbolic acid, 431. Carbon, 181. Amorphous, 181, 184, 190. Boride, 261. Combining power, 202. Compounds, 181, 405, 406. Bisulphide, 112, 252. Gas, 190, 213. Silicide, 117, 258. Test, 189. Carbonado, 182. Carbonate, acid, 195. Ammonium, 152. Carbonates, 195. Normal, 195. Carbon dioxide, 191. And combustion, 193. Composition, 195. .Detection, 67, 68. Formation, 191. History, 196. In air, 194. In atmosphere, 67. Liquid, 193. Occurrence, 191. Other names, 203. Preparation, 192. Properties, 193. Relation to life, 194. Solid, 193. Solubility, 42, 193, 194. Test, 192, 325. Carbonic acid, 97, 194. Anhydride, 195. Oxide, 197. Carbon monoxide, 196, 197. In water gas, 215. Carbonyl chloride, 198. Nickel, 198. Carborundum, 117. Furnace, 117, 118. Carboxyl, 414. Carbureter, 213. Carbureting, 215. Carlisle and Nicholson, 53, Carnallite, 294, 295, 331. Magnesium from, 332. Carnelian, 255. Cassiterite, 354. Casting iron, 378. Cast iron, 378. Varieties, 378. Castner, 284, 285. Catalysis, 250. Catalytic action, 250. Catalyzer, 250. Cathode, 120, 121. Cation, 120, 121. Caustic lime, 324. Lunar, 312. Potash, 297. Soda, 290. Cavendish, 16, 27, 30, 55, 64, 69, 157, 443- Celestite, 328. Cell, electrolytic, 122. Galvanic, 119. Voltaic, 119. Celluloid, 429. Cellulose, 428. Nitrates, 428. Cementation process, 380. Cements, 325. Centigrade thermometer, 439. Cerussite, 357, 361. Chalcedony, 255. Chalcopyrite, 301, 373. Chalcorite, 301. Chalk, 322, 323. Chalybeate water, 37, 387. Index. 563 Champagne, 193. Changes, I. Chaptal, 72. Charcoal, 187. Animal, 187. Pit, 1 88. Wood, 187. Charles, 444. Law, 19. Checkerberry, 432. Chemical action, 3, ill, 250. And electricity, 119. And heat, 112. And light, 51, ill. And solution, 47. And temperature, 113. Classes, 3. Chemical attraction, 4. Chemical calculations, 103. Chemical change, I, 2, 14, 47. And ozone, 22. Chemical compounds, 69. Chemical energy, in. Chemical equivalents, 123. Chemicking, 138. Chemism, 4. Chemistry, defined, I, 2. Organic, 405. Chemists' table, 447. Chest, 256. Chili saltpeter, 231, 292. Chinese white, 336. Chloral, 412. Hydrate, 412. Chloride, of lime, 137. Test, 144. Chlorides, 139, 140, 143, 144. Chlorination process, 315. Chlorine, 133. Acids, 91. And hydrogen, 136. And water, 57. Available, 137. Compounds, 296. Determination of atomic weight, 171 . Chlorine, continued. Hydrate, 139. Liquid, 139. Name, 135. Nascent, 139. Occurrence, 133. Preparation, 133-134. Properties, 135. Uses, 139. Water, 51, 135. Chloroform, 412. Chlorophyll, 373. Chloroplatinic acid, 394. Choke damp, 203. Chroma. 365. Chromates, 366, 368. Chrome alum, 350, 367, 368. Iron ore, 365. Orange, 368. Red, 368. Steel, 366. Yellow, 367. Chromic chloride, 368. Compounds, 368. Hydroxides, 368. Oxide, 368. Sulphate, 368. Chromite, 365. Chromites, 368. Chromium, 365. As a metal, 368. Compounds, 366, 368. In minerals and rocks, 365. Name, 365. Ore, 365. Silicide, 258. Tests, 367, 368. Trioxide, 368. Uses, 365. Chromous compounds, 368. Chronological table of chemists, 447. Cinchona tree, 433. Cinchonine, 433. Cinder, 375. Cinnabar, 338, 5 6 4 Index. Citric acid, 419. Classification, organic compounds, 408. Periodic, 398. Clay, 351, 352. Aluminium from, 345. Clouds, 67. Coal, 184. And graphite, 118. Beds, 184, 185. Bituminous, 185, 186, 210. Composition, 186. Distillation, 210. Distribution, 187. Fields, 1 86, 187. Fire, 196. Gas, 210. Gas plant, 211. Mines, gases in, 202, 203. Products from, 213. Section, 185. Soft, 185, 189. Coal tar, 213. Dyes, 431. Cobalt, 389. Blue, 390. Test, 390. Cocaine, 433. Coca plant, 433. Coffee, 433. Coins, gold, 316. Nickel, 306, 389. Silver, 311. Coke, 189, 213. As fuel, 190. Coal gas from, 210. From petroleum, 209. In iron smelting, 377. Colemanite, 262. Collodion, 428. Colored glass, 260. Color of metals, 279. Combination, 3. By volume, 53. By weight, 53. Of gases, 1 66. Combustion, 15, 1 6, 28, 67, 148, 191. And flame, 217. Old theory, 15. Products, 218. Spontaneous, 14. Common salt, 133, 286, 287. See So- dium chloride. Complete fertilizer, 271. Components, 8. Composition, ammonia gas, 153. Coal, 1 86. Carbon dioxide, 195. Earth's crust, 6. Heavenly bodies, 404. Hydrochloric acid, 143. In per cent, 103. Natural waters, 38. Nitric acid, 157. Nitric oxide, 162. Nitrous oxide, 161. Of a compound, 50. Organic compounds, 405. Water, 2^ 27, 57. Compounds, chemical, 8, 69. Saturated, 177. Unsaturated, 177, 178. Concentrated, defined, 41. Alum, 349. Concentration of ore, 281. Condenser, 39, 40. Coal gas, 212. Iodine, 231. Conductivity, metals, 279. Solutions, 126-127. Cones, Bunsen flame, 220. Flame, 216-217. Conservation, energy, in. Matter, 4, 104. Constitution, benzene, 430. . Organic compounds, 406. See Composition. Constitutional formula, 407. Contact method for sulphuric acid, 24^ Converter, 381. Cooking soda, 290. Index. 565 Copper, 301. Acetate, 417. Alloys, 305, 316. And sulphuric acid, 243. Arsenite, 243. Carbonates, 301, 302, 303. Coins, 306. Compounds, 306. Electrolytic, 303. Fluoride, 227. From Michigan, 301, 302. Glance, 301. History, 301. Iron sulphides, 301, 302, 373. Metallurgy, 302. Name, 301. Native, 301, 302. Nitrate, 307. Ores, 301. Oxides, 301, 302, 303, 306, 307. Production, 302. Properties, 303. Purification, 303. Pyrites, 301. See Cupric and Cuprous. Smelting, 302. Region, map, 374. Replacement, 304. Replacing power, 304. Silicide, 258. Sulphate, 307. Sulphide, 301, 307. Test, 304. Uses, 304. Copperas, 385. Coquina, 322. Coral, 323. Cordials, 411. Corrosive sublimate, 340. Corundum, 343, 346, 347. Cottolene, 421. Courtois, 231, 444. Crayon, 323. Cream of tartar, 290, 418. Potassium carbonate from, 297. Crockery, 352. Crocoisite, 365. Crocoite, 365. Crocus, 384. Crucible process for steel, 380. Cruikshank, 119. Cryolite, 225, 343, 344, 350. Crystal, rock, 255. Crystals, 44, 440. Hexagonal, 441. Isometric, 441. Monoclinic, 442. Orthorhombic, 239, 441. Production, 440. Snow, 35. Systems, 440. Tetragonal, 441. Triclinic, 442. Crystallization, 44, 440. Water of, 45, 46. Cubic cleavage, 363. Cupel, 310. Cupellation, 310, 360. Cupric compounds, 306. Oxide, 307. Sulphate, 307. Sulphide, 307. Cuprite, 301, 306. Cuprium aes, 301. Cuprous compounds, 306. Oxide, 306, 426. Sulphide, 307. Cuprum, 301. Current, electric, 121. Cyanic acid, 198. Cyanide, mercury, 198. Potassium, 198. Process, 198, 315. Iron, 387. Cyanogen, 198. Cymogene, 208. Cyprian brass, 301. Dalton, 77, 79, 444. Davy, 53, 97, 119, 135, 161, 182, 566 Index. 221, 231, 284, 286, 319, 343, 444- Deacon's process for chlorine, 134. Decay, 17, 67, 155, 191. Decomposition, 3. Double, 3, 45. Heat of, 1 1 3. Definite proportions, law, 76. Deflagration, 159. Dehydrated, 46. Deliquescence, 46, 67. Destructive distillation, 188, 202. Determination, atomic weights, 170- 171. Developer, 313. Deville, 343. Dewar, 29, 444. Bulb, 70. Dew point, 66. Dextrin, 427. Dextrose, 425, 426. Diacid base, 94. Diamond, 181, 182, 190. Artificial, 182. Cheap, 257. Drill, 182. Diatomaceous earth, 256, 257. Diatoms, 256. Dibasic acid, 92. Dicalcium phosphate, 271. Dichlorethane, 204. Dichromates, 368. Diffusion, 26, 68. Diluents, 215. Dilute, 41. Diphosphates, 269, 271. Disinfectant, carbolic acid, 431. Formaldehyde, 413. Disodium phosphate, 269. Displacement, downward, 135. Upward, 148. Dissociation, by heat, 151. Electrolytic, 125, 127. Distillate, 40. Distillation, 39. Distillation, continued. Coal, 210. Destructive, 188, 202, 204. Dry, 147. Petroleum, 208. Water, 39. Wood, 1 88, 409. Distilled liquors, 411. Water, 40. Dolomite, 331, 334. Double decomposition, 3, 45. Refraction, 320. Downward displacement, 135. Drinking water, 39. Lead in, 359. Drummond light, 29. Ductile metals, 279. Dulong, 444. And Petit, 172. Dumas, 182, 397, 444. And Boussingault, 66. And Stas, 56, 57. Dutch leaf, 305. Metal, 305. Process for white lead, 361. Dyads, 176. Dyeing, 350. Dynamite, 422. Earthenware, 352. Effervescence, 42, 193. Effervescing powder, 290. Efflorescence, 46. Electrical conductivity, 126^ Electric battery, 120. Electric furnace, 114-115, 184, 365. Industrial use of, 116. Electricity and chemical action, 119. Electric light carbons, 209. Electrochemical equivalent, 123. Terms, 120. Electrochemistry, 119. Electrodes, 118, 120, 121, 184, 190. Electrolysis, 120. Aluminium oxide, 343. Index. 56? Electrolysis, continued. And solution, 126. Calcium iodide, 319. Carnallite, 332. Copper sulphite, 303. Galena, 358. Gold solution, 315. Hydroxides, 284. Illustrations, 122. Industrial application, 124. Metals, 281. Potassium hydroxide, 294. Sodium chloride, 122, 291. Sodium hydroxide, 284. Sodium nitrate, 302. Theory of, 125. Water, 52, 123. Zinc chloride, 122. Electrolytic cell, 122. Copper, 303. Dissociation, 125. Process for chlorine, 134. Process for white lead, 362. Separation of gold and silver, 315. Electro-negative ions, 121, 122. Positive ions, 121, 122. Silicon, 257. Thermal manufacture of carbon disulphide, 252. Electrolyte, 120, 128. Electroplating, 1 24- 125. Electrotyping, 124-125. Elements, 5, 6, 7, 448, 449. Acid properties, 396. Basic properties, 396. Bivalent, 176. Classification, 396. Families, 397. General relations, 396. In earth's crust, 6. In organic compounds, 405. In sun, 404. Numerical relations, 397. Periodic classification, 398. Prediction, 401. Elements, continued. Quadrivalent, 176. Quinquivalent, 176. Spectra, 402. Trivalent, 176. Table, 448, 449. Univalent, 176. Emerald, 347. Emery, 343, 346. Empirical formula, 178, 407. Emulsin, 432. Endothermic, 112. Energy, chemical, 4. Mechanical, 33. Enriching gas, 213, 215. Epsom salts, 333. Equation, 83, 84. Gas, 175. Illustrating reactions, 106. Ionic, 129, 130. Molecular, 175. Problems based on, 107. Quantitative significance, 104. Thermal, 112. Volumetric, 175. Equivalents, 100. And valence, 178. Chemical, 123. Electrochemical, 123. Multiples, 101. Table, 100. Erosion, 32. Esters, 419. Etching, 227. Ethane, 202, 203, 409. Graphic formula, 407. Ether, ethyl, 413. And water, 43. Sulphuric, 414. Ethereal salts, 419, 420. Ethers, 413. Ethyl, 406, 409. Acetate, 419. Alcohol, 406, 408, 409. Butyrate, 419. 568 Index. Ethyl, continued. Ether, 413. Oxide, 414. Sulphuric acid, 414. Ethylene, 204. Chloride, 204. In illuminating gas, 216. Series, 202. Eudiometer, 53, 54. Evaporation, 440. Exercises, 9, 20, 30, 48, 58, 73, 85, 98, 108, 130, 145, 163, 178, 198, 222, 233, 253, 263, 276, 282, 299, 3*7 3 2 9, 340, 35 2 3 6 3 372, 390, 394, 404, 433- Exhauster, 212. Exothermic, 112. Exposure, photographic, 313. Factors, 83. Fahrenheit thermometer, 439. Families of elements, 397. Faraday, 120, 123, 128, 139, 193, 444. Law, 123. Fats, 420. Fatty acid series, 414. Fehling's solution, 426. Feldspar, 293, 343, 351. Fermentation, 192, 410. Acetic, 416. Alcoholic, 410. In bread making, 421. Sugar, 410. Ferments, 410. And glucosides, 432. Ferric compounds, 384. Chloride, 386. Ferrocyanide, 388. Hydroxides, 385. Oxide, 384. Sulphate, 385. Sulphide, 386. Ferricyanides, 387, 388. Ferrochrome, 366. Ferrocyanides, 387, 388. Ferro-ferric oxide, 385. Ferromanganese, 369, 378. Ferrous compounds, 384. Carbonate, 387. Chloride, 386. Ferric oxide, 384. Ferricyanide, 388. Ferrocyanide, 388. Hydroxide, 385. Sulphate, 385. Sulphide, 240, 385. Ferrum, 373. Fertilizer, 73, 271. Manufacture, 271. Potassium salts as, 298. Sodium nitrate as, 292. Film, photographic, 313. Filter, charcoal, 188. Filtering water, 39. Fire, 15. Damp, 202. Extinguisher, 194. Fireworks, 14, 332. Fixed air, 196. Alkalies, 93. Fixing, in photography, 303. Flame, 216. Acetylene, 207, 216. And combustion, 217. Bunsen, 219. Hydrogen, 27, 112. Non -luminous, 219, 220. Oxidizing, 221, 222. Oxyhydrogen, 29. Parts, 216-217. Reducing, 222. Smoky, 218. Flashing point, 209. Flavors, 419. Flint, 255, 256. Flour, wheat, 427. Flowerpots, 352. Flowers of sulphur, 238. Fluid, magnesia, 334. Fluorides, 227, 343. Index. 569 Fluorine, 225. Apparatus, 226. Isolation, 225. Liquid, 226. Name, 225. Properties, 226. Fluor spar, 225, 226. Flux, 281, 375. Food, water in, 31, 32. Fool's gold, 386. Formaldehyde, 412. Formalin, 413. Formation, heat of, 1 12. Formula, 82. Constitutional, 407. Empirical, 178,407. Graphic, 178, 407, 413, 414. Molecular, 174. Rational, 407. Simplest, 104, 174, 175. Structural, 178, 407, 413, 414. Fossil, from coal bed, 185. Frame, for soap, 423. Franklinite, 334. French process for white lead, 362. Fructose, 425. Fruit sugar, 425. Fuming acid, nitric, 163. Sulphuric, 251. Furnace, blast, 281, 375. Reverberatory, 281, 282. Fusible alloys, 337, 360. Metals, 275. Fusion, for crystals, 440. Gahnite, 334. Galena, 357, 362. Crystals, 362. Gallic acid, 432. Gallium, 401. Galvanic cell, 119. Galvanized iron, 336. Gangue, 280. Gaps in periodic system, 401. Garnet, 347. Gas, 61. Carbon, 190, 213. Coal, 210. Effect of heat on volume, 18, 19. Effect of pressure on volume, 18. Equation, 175. Flame, structure, 218. Holder, 212. Illuminating, 210. Marsh, 202. Natural, 209. Producer, 25. Sylvestre, 196. Volume, reduction, 53, 54. Water, 25, 196, 213. Water, plant, 214. Gases, absorption by charcoal, 188. By platinum, 394. Combination by volume, 166. Inert, 69. In mines, 221. Properties, 166. Solution of, 41. Gasolene, 208. Gay-Lussac, 55, 231. Law, 1 66. Tower, 249. Gelatine plate and film, 313. Gems, aluminium, 347. Artificial, 347. Glass, 260. Quartz, 257. Generator, acetylene, 207. Water gas, 213. German process for white lead, 362. Silver, 306, 389. Geyserite, 258. Gin, 411. Glacial acid, acetic, 415. Phosphoric, 268. Glass, 258. And hydrofluoric acid, 227. Annealing, 260. Blasting, 257. Blowing, 259, 260. 570 Index. Glass, Bohemian, 260} Colored, 260. Constituents, 259. Crown, 260. Cut, 260. Flint, 260. Kinds, 258, 259. Manufacture, 259. Plate, 259. Polishing, 260. Production, 260. Typical mixture, 259. Window, 259. Glauber, 140, 445. Salt, 292. Glazing pottery, 352. Globigerina ooze, 322, 323. Glover tower, 248. Glucose, 425, 426. Glucosides, 432. Glycerides, 420. Glycerine, 420. Preparation, 422. Properties, 421. Relation to soap, 420. Uses, 421. Glycerol, 422. Glyceryl, 429. Oleate, 420. Palmitate, 420. Stearate, 420. Gneiss, 255. Gogebic iron range, 374. Gold, 314. Alloys, 314, 316. Amalgam, 339. Chloride, 315, 316, 317. Coin, 316. Compounds, 317. Cyanide, 315, 317. Distribution, 309, 314. Dust, 314. Dutch, 305. Finely divided, 317. Gold, continued. Fool's, 386. History, 313. Leaf, 316. Making, 314. Map of distribution, 309. Name, 314. Nugget, 314. Parting, 315. Pen tips, 394. Plating, 317. Production, 314. Properties, 316. Purification, 315. Red, 316. Reduction of compounds, 317. Separation from silver, 315. Test, 317. Uses, 316. White, 316. Graham, 26, 445. Gram, 437. Granite, 255. Grape sugar, 425, 426. Graphic formula, 178, 407, 413, 414. Graphite, 183, 190. Artificial, 118. Gravimetric, 53. Composition, air, 66. Composition, water, 55, 57. Gray cast iron, 378. Green fire, 329. Pigments, 368. Vitriol, 246, 385. Grindstones, 256. Groups of elements, 397. Guano, 271, 331. Guignet's green, 368. Gun cotton, 428. Metal, 305. Gunpowder, 14, 296. Smokeless, 428. Gypsum, 326. Reduction of, 235. Index. Haemoglobin, 373. Halides, 96, 225. Hall, 343- Process for aluminium, 343, 344. Halogens, 225. Haloid salts, 225. Hardness, of metals, 279. Of water, 37, 327. Permanent, 327. Temporary, 327. Hard water, 37. Coal, 185, 1 86. Harveyized steel, 380. Hausmannite, 369. Heat, and chemical action, 112, H3- And oxidation, 14. From burning hydrogen, 1 12. In electric furnace, 114. Of decomposition, 113. Of formation, 112. Of neutralization, 130. Heavenly bodies, constitution, 404. Helium, 69, 404. Hematite, 373. Henry's law, 42. Heroult process for aluminium, 344. Hexagonal crystals, 441. Hofmann, 445. Apparatus, 52. Honey, 425. Horn silver, 133, 308. Humboldt, 55. Hydrargyrum, 338. Hydrate, 93. Chlorine, 139. Hydrated, 46. Hydraulic lime, 325. Main, 210. Mining, 314. Hydriodic acid, 232. Hydrocarbons, 202, 408. Hydrobromic acid, 230. Hydrochloric acid, 140-143, Commercial, 141. Hydrochloric, continued. Composition, 143. Liquefied, 142. Test, 144. Hydrocyanic acid, 198. Hydrofluoric acid, 227, 257. Vapor density, 228. Hydrogen, 23. And chlorine, 136. And periodic classification, 401. And steam, 24. And water, 50. Arsenide, 273. Chemical conduct, 27. Diffusion, 26. Dioxide, 59. Discovery, 30. Explosions, 28. Flame, 27. In acids, 24, 87, 90. Ions, 121. Liquid, 29. Name, 25, 30. Peroxide, 59. Physical properties, 25. Preparation, 24. Solid, 29. Valence, 176. Weight of liter, 25. + Hydrogen sulphide, 240, 241, 242. Composition, 241. .-- Test, 242. Water, 241. Hydroquinone, 431. Hydroxides, 89, 93. And alcohols, 409. Common names, 93. Organic, 409. Hydroxyl, 89, 94. Hygroscopic, 46. Hypo, 91, 252. Hypophosphites, 269. Hyposulphite in photography, 313. Hypothesis, 76. Avogadro's, 167. 572 Index. Ice, 32, 34, 35. Making plant, 153. Manufactured, 153. Stone, 350. Iceland spar, 320. Illuminants, 216. Illuminating gas, 210. Carbon monoxide in, 197. Characteristics, 215. Composition, 215. Illuminating power, 216. Impurities, 240. Luminosity, 216. Indicator, 98. Inert gases in atmosphere, 69. Infusorial earth, 256, 257. Ingots, 381. Ink, 385, 418. Indelible, 312. Printer's, 190. Writing, 433. Inorganic compounds, 405. Insoluble substances, 41. Sulphate, test, 251. Intervals in periodic classification, 398. Iodides, 232. Iodine, 230. Commercial preparation, 231. Compounds, 232. Detection, 232. Determination, 252. Discovery, 231. In seaweed, 230. Name, 232. Preparation, 230. Production, 233. Properties, 231. Purification, 231. Source, 293. Test, 232. Uses, 233. Vapor density, 232. lodoform, 233, 412. Ionic equation, 129, 130. lonization, 125. And acids, bases, and salts, 129. Application, 129. Table, 127. Ions, 120, 121, 125, 126. Test for, 129. Iridium, 226, 392, 393, 394. Iridosmine, 394. Iron, 373. Acetate, 417. Alum, 386. And coke, 190. By alcohol, 383. By hydrogen, 383. Carbide, 209, 285. Carbonate ores, 374. Cast, 377, 378. Chemistry of smelting, 377. Chlorides, 386. Compounds, 384. See Ferric and Ferrous. Cyanides, 387. Disulphide, 386. Galvanized, 386. History, 373. Impurities, 377. Liquor, 417. Magnetic oxide, 385. Malleable, 879. Map of deposits, 374. Metallurgy, 379. Ore, 373, 374. Ore, chrome, 365. Ore, consumption, 377. Ore, deposits, 374. Ore, reduction, 197, 375. Oxides, 384. Passive, 384. Pig, 377- Properties, 383. Pyrites, 373, 385, 386. Rust, 383. Rusting, 14. Silicide, 258. Smelting, 375. Index. 573 Iron, continued. Spiegel, 369. Sulphides, 386. Symbol, 373. Test, 388. Varieties, 377. Isomerism, 204. Isomers, 204. Isometric crystals, 441. Ivory black, 189. Jasper, 256. Javelle's water, 139. Kainite, 294, 298, 331. As fertilizer, 298. Kali, 204. Kalium, 294. Kaolin, 351, 352. Kassiteros,, 354. Kelp, 231. Kerosene, 209. Kieserite, 331, 333. Kilogram, 437. Kindling temperature, 113,218, 221. Kirchhoff, 403, 445. Krypton, 69, 404. Labarraque's solution, 139. Lactic acid, 290, 418. Lactose, 425. Lake, 350. Lampblack, 190. Laudanum, 433. Lavoisier, 5, 15, 16, 18, 25, 27, 30, 50, 55, 63, 64, 88, 97, 157, 182, 196, 396, 445. Law, 75. Boyle, 19. Charles, 19. Conservation of energy, in. Definite proportions, 75, 76, 79. Faraday, 123. Gay-Lussac, 166. Henry, 42. Law, continued. Matter, 5. Multiple proportions, 77, 78. Periodic, 398. Specific heat, 172. Lead, 357. Acetate, 363,417. Alloys, 360. Argentiferous, 308. Black, 183, 359. Carbonate, 357, 361. Carbonate, basic, 361. Chambers, 249. Chloride, 363. Chromate, 367. Chromate, native, 365. Compounds, 363. Compounds, poisonous, 359. Cupellation process, 310. Dioxide, 361. History, 357. Hydroxide, 362. In drinking water, 359. Interaction with metals, 359. Metallurgy, 358. Monoxide, 360. Nitrate, 363. Nitrate, behavior with heat, 163. Ore, 357. Oxides, 360. Parkes process for, 309. Pencils, 184. Peroxide, 361. Phosphate, 357. Pipe, 366. Production, 357. Properties, 358. Silver bearing, 308. Spongy, 360. Sugar of, 363, 417. Sulphate, 357, 363. Sulphide, 242, 357, 362, 363. Test, 363, Tetroxide, 360. 574 Index. Lead, continued. Uses, 359. White, 361. Leather, 433. Leblanc process for sodium carbonate, 288. Lemon juice, 90. Levulose, 425. Liebig, 230, 445. Life and carbon dioxide, 194. Oxygen, 16. Nitrogen, 72. Phosphorus, 270. Potassium, 298. Light and chemical action, 51, ill. Silver salts, 312,313. Lignite, 185. Lime, 324. Air slaked, 324. And water, 113, 324. Caustic, 324. Chloride of, 137. Hydraulic, 325. Light, 29, 324. Making, 192, 324, 325. Milk of, 326. Quick, 324. Superphosphate, 271. Uses, 324. See Calcium oxide. Limekiln, 193, 325. Limestone, 320. As flux, 377. Burning, 325. Caves, 321, 322. Fossil, 322. Solubility, 321. Uses, 323. Lime water, 325. And carbon dioxide, 192, 325. Detection, 68. Preparation, 326. See Calcium hydroxide. Liming, 138. Limonite, 373. Links, 407. Liquid air, 12, 69. Acetylene, 205. Ammonia, 148-149, 153. Carbon dioxide, 193. Chlorine, 139. Fluorine, 226. Hydrogen, 29. Oxygen, 18. Sulphur dioxide, 244. Liquids, solubility, 43. Liquor, alcoholic, 411. Distilled, 411. Iron, 417. Red, 350. List of reference books, 450. Litharge, 360. Lithia water, 298. Lithium, 298. Citrate, 298. Discovery, 294. Test, 298. Litmus, action on, acid, 90. Alkali, 92. Base, 92. Neutral substance, 94. Salt, 94. Loadstone, 385. Lubricating oil, 209. Luminosity, illuminating gas, 2 1 6. Of flame, 218. Luminous paint, 329. Lunar caustic, 312. Luray cavern, 321, 322. Luster, 279. Madder, 432. Magnalium, 346. Magnesia, 333, 334, 370. Alba, 334, 370. Black, 370. Fluid, 334. Mixture, 333. Nigra, 370. Index. 575 Magnesia, continued. Stone, 370. Uses, 333. Magnesite, 334. Magnesium, 331. Alloy, 346. Bromide, 228. Calcium carbonate, 331. Carbonate, 331, 334. Chloride, 333. Citrate, 334. Compounds in soil, 331. Compounds and water, 327. Hydroxide, 333. Nitride, 153,332. Oxide, 333. See Magnesia. Phosphates, 331. Preparation, 332. Properties, 332. Ribbon, 332. Sulphate, 333. Uses, 332. Magnetic oxide of iron, 385. Magnetite, 373, 385. Majolica, 352. Malachite, 301, 308. Malic acid, 418. Malleable iron, 379. Metals, 279. Mammoth cave, 322. Manganates, 371. Manganese, 369. Alloys, 369. As non-metal, 371. Black oxide, 370. Compounds, 371. Dioxide, 369. Isolation, 370. History, 370. Name, 370. Ores, 369. Preparation, 369. Production, 369. Properties, 369. Test, 372. Manganese, continued. Uses, 369. Manganesium, 370. Manganite, 369. Manganous compounds, 371. Chloride, 370, 371. Hydroxide, 370. Sulphate, 371. Sulphide, 371. Mantle, Welsbach, 222. Map, copper deposits, 374. Gold, 309. Iron, 374. Silver, 309. Marble, 320. Marchand tube, 56. Marengo cave, 322. Marquette iron range, 374. Marsh gas, 202. Marsh's test for arsenic, 273. Massicot, 360. Matches, 270. Matte, copper, 302. Matter, conservation, 4. Properties, i, 2. Meadowsweet, 430. Meerschaum, 331. Mendeleeff, 398, 445. Menominee iron range, 374. Mercuric chloride, 340, 357. Cyanide, 198. Nitrate, 340. Oxide, 1 8, 339. Sulphide, 338, 340. Mercurous chloride, 339, 357. Nitrate, 340. Mercury, 337. Alloys, 339. Compounds, 339. Deposits, 338. Fulminating, 339. Name, 338. Native, 337. Ore, 338. Preparation, 338, 57 6 Index. Mercury, contintced. Production, 338. Properties, 338. Purification, 338. Specific heat, 172. Transportation, 338. Uses, 339. Vapor density, 169, 339. Mesabi iron range, 374. Metal, and non-metal, 278. Babbit's, 336. Bath, 305. Bell, 306. Britannia, 306, 356. Dutch, 305. Gun, 305. Hypothetical, 150. Muntz, 305. Newton's, 275. Rose's, 275. Speculum, 306. Type, 360. White, 306. Wood's, 275, 337. Metallic ions, 121. Luster, 279. Metalloids, 278. Metallurgy, 280. Copper, 302. Lead, 358. Iron, 375. Silver, 309, 310. Metals, action with nitric acid, 158. Alkali, 284. Alkaline earth, 319. Antifriction, 336. Chemical properties, 279. Classification, 396. Familiar, 7. Found free, 280. General properties, 278. Known to ancients, 280. Occurrence, 279. Physical properties, 278. Platinum, 394. Metals, continued. Preliminary treatment, 280. Preparation, 280. Metamerism, 204. Metaphosphates, 269. Metaphosphoric acid, 268. Metastannic acid, 355. Metathesis, 3. Meter, defined, 437. Gas, 212. Methane, 202, 409. Graphic formula, 407. In natural gas, 209. Series, 202. Methyl, 406, 409. Alcohol, 409. Benzene, 430. Salicylate, 432. Methylated spirit, 410. Metric abbreviations, 438. Apparatus, 394. Equivalents, 438. System, 437. Ton, 32. Transformations, 438. Mexican onyx, 322. Meyer, Lothar, 398, 445. Mica, 293, 343. Microcosmic salt, 269. Milk of lime, 326. Sulphur, 240. Milner's process for white lead, 362. Mineral, defined, 280. Compounds, 405. Springs, 37, 42. Water, 37. Minerals, 258. Minium, 360. Mispickel, 272. Mixture, 9, 77. Air, 69. Modification, allotropic, 191. Moissan, 114, 116, 182, 184, 225, 226, 3!9. 365, 445- Moissan's electric furnace, 1 14. Index. 577 Molecular equation, 175. Formula, 174. Molecular weights, 103, 128, 168. And vapor density, 168. Determination, 170, 171. Exact, 170. Hydrogen, 169. Steam, 169. Molecules, 80-8 1, 167-168. And atoms, 80. And equations, 175. Molybdenum, 369. Monacid base, 94. Monads, 176. Monobasic acids, 92. Monocalcium phosphate, 271. Monoclinic crystals, 442. Sulphur, 239. Monophosphates, 269. Mordants, 350, 357, 367. Morphine, 433. Mortar, 326. Moth balls, 432. Mother liquor, 230, 231. Mucilage, 427. Multiple proportions, law, 77-78. Table, 78. Muntz metal, 305. Muria, 140. Muriate of ammonia, 151. Muriatic acid, 92, 140. Muscovado sugar, 424. Mutton fat, 420. Naphtha, 208. Naphthalene, 432. Nascent state, 138. Natrium, 284. Natron, 284. Natural gas, 209. Natural groups, 400. Waters, 38. Nature of solution, 48. Negative electrode, 121. Photographic, 313. Neon, 69, 404. Neutral, 94. Reaction, 94. Neutralization, 88, 89, 97. And ionic theory, 130. Heat of, 130. Newton's metal, 275. Niagara Falls, industries at, 1 1 6, 117, 118, 155, 291, 344. Nicholson and Carlisle, 53, 119. Nickel, 388. Alloys, 306. Carbonyl, 198. Coin, 306, 389. Hydroxide, 389. Ores, 388. Plating, 389. Properties, 389. Steel, 383, 389. Test, 389. Uses, 389. Nickeloid, 389. Nicotine, 433. Niter, 72. Meal, 295. Source, 295. Nitrates, 158. Behavior with heat, 159. Deposits, 155. Test, 159. Nitric acid, 154, 155, 156. Action with metals, 158. And copper, 159, 162. And electric sparks, 155. Composition, 157. Formation, 155. Fuming, 163. Preparation, 155. Test, 159. Uses, 157. Nitric oxide, 159, 162. Composition, 162. Nitrides, 72. Magnesium, 153. Nitrification, 155. 578 Index. Nitrites, 159. Nitrogen, 72. Discovery, 63. Effect on flame, 220. In atmosphere, 63. Name, 72. Oxides, 78, 1 60. Pentoxide, 163. Peroxide, 159, 162, 163. Preparation, 72. Properties, 63, 72. Proportion in air, 64. Relation to life, 72. Tetroxide, 163. Trioxide, 163. Valence, 177, 178. Nitrous acid, 159. Nitrous oxide, 160, 161. Composition, 161. Discovery, 161. Nitrobenzene, 430. Nitroglycerine, 422. Nitrosyl-sulphuric acid, 248. Nomenclature, acids, 91. Bases, 93. Hydroxides, 93. Salts, 95. Non-luminous flame, 219, 220. Non-metallic ions, 121. Non-metals, 88. Classification, 396. General properties, 278. Nordhausen sulphuric acid, 252. Normal bismuth nitrate, 276. Normal salts, 96. Nugget, gold, 314. Occlusion, 26. Ocean water, 38. Salts in, 38. Oil, and water, 43. Lamp flame, 218. Lubricating, 209. Of bitter almonds, 431, 432. Of vitriol, 92, 246. Oils, 420. Olefiant gas, 204. Olein, 420. Oleomargarine, 421. Olive oil, 420, 421. Onyx, 255. Opal, 256. Opaque, 279. Open hearth process for steel, 382. Opium, 433. Orange mineral, 361. Ore, defined, 280. Calcination, 281. Classes, 280. Dressing, 281. Organic acids, 92, 414. Chemistry, 405. Compounds, 405, 406, 408. Orpiment, 272, 273. Orthophosphoric acid, 268. Orthorhombic crystals, 441. Sulphur, 239. Osmium, 392, 394. Ostwald, 445. Oxalic acid, 417. Oxidation, 14, 192, 357. And decay, 17. By potassium permanganate, 371, Of food, 1 6. Oxide, carbonic, 197. Oxides, 15. Acidic, 97. Basic, 97. Of nitrogen, 160, 246-248. Relation to acids and bases, 96. Oxidized silver, 311. Oxidizing agent, 14, 60. In matches, 270. Oxidizing flame, 221, 222. Oxychloride, antimony, 275. Bismuth, 276. Oxygen, 11. Absorption by silver, 311. And blood, 16, 17. And combustion, 15. Index. 579 Oxygen, continued. And flames, 218. And ozone, 22. And water, 51. Breathing pure, 17. Erin's process, 12. Discovery, 18. In acids, 87, 88, 91, 97. In atmosphere, 63. Liquid, 1 8. Name, 18, 88. Nascent, 138. Preparation, II, 293. Properties, 12. Relation to life, 16. Solid, 1 8. Uses, 17. Weight of liter, 18. Oxyhydrogen blowpipe, 17, 28, 29. Oxymuriate, tin, 357. Ozone, 21, 113. In atmosphere, 62. Formula, 169. Paint, black, 190. Blue, 417. Lead, 357. Luminous, 329. Red, 273, 340, 361, 384. White, 336, 362. Yellow, 273, 367. Pakfong, 306. Paktong, 306. Palladium, 392, 394. Absorption by, 26, 394. Palmitic acid, 417. Palmitin, 420. Palm oil, 417. Paper, making, 429. Parchment, 428. Paracelsus, 30. Paraffin, series, 203. Wax, 209. Paregoric, 433. Pads green, 273, 417. Parkes process for silver, 309. Parting, gold and silver, 315. Passive iron, 384. Paste, gems, 262. Glass, 260. Starch, 427. Pastry, raising, 290. Pearlash, 297. Peat, 185. Pentads, 176. Percentage composition, 103. Periodic classification, 398. Gaps, 401. Periodic law, 398. Periodic process for bromine, 229. Periodic table of elements, 399. Periods in periodic classification, 398. Permanent hardness, 327. Peroxide, hydrogen, 59. Sodium, 293. Petit, 172,445. Petrified wood, 256, 257, 258. Petroleum, 207-209. Origin, 209. Production, 209. Refining, 208. Pewter, 356, 360. Phenol, 431. Derivatives, 431. Phenyl, 406. Methane, 406. Philosopher's stone, 314. Phlogiston, 15, 1 8. Phosgene, 198. Phosphates, 265, 269. Acid, 26$. Dicalcium, 271. Disodium, 269. Monocalcium, 271. Primary, 269. Rock, 271. Secondary, 269. Slag, 271. Tricalcium, 271. Phosphine, 269. 5 8 Index. Phosphonium compounds, 269. Phosphor bronze, 305. Phosphoric acids, 268. Oxide, 268. Phosphorite, 265. Phosphorous oxide, 268. Phosphorus, 265. Acids, 268. Action on air, 65, 72. And ozone, 21. And plants, 270. Black, 267. Discovery, 265. Electrolytic manufacture, 266. In plants and animals, 265. Manufacture, 265, 266. Minor compounds, 269. Name, 267. Ordinary, 266. Oxides, 268. Pentachloride, 270. Pentoxide, 65, 268. Properties, 266. Purification, 266. Red, 267. Relation to life, 270. Salts, 268. Trichloride, 270. Uses, 267. Vapor density, 169, 267. Yellow, 266. Photography, ill, 312. Photometer, 216. Phylloxera, 240. Physical changes, I, 2. Pickles, 90, 416. Picrates, 431. Picric acid, 431. Picromerite, 294. Pig iron, 377. Pinchbeck, 305. Placer mining, 314. Plants and atmosphere, 194. And nitrogen, 72, 73. And phosphorus, 270. Plants and atmosphere, continued. And potassium, 298. And silica, 257. Plaster, 326. Of Paris, 327. Plata, 392. Plate, developing, 313. Photographic, 312. Platina, 392. Platinic chloride, 394. Platinum, 392. Absorption of gases, 394. Alloys, 394. And aqua regia, 392. And iridium, 226, 392. And sulphur dioxide, 245. And sulphuric acid, 249. Arsenide, 392. Black, 394. Compounds, 394. Discovery, 392. Dish, 393. Foil, 393. In electric light bulbs, 393. Metals, 394, 401. Name, 392. Native, 392. Ore, 392. Preparation, 392. Print, 394. Production, 392. Properties, 393. Sheet, 393. Source, 392. Spongy, 392, 393. Uses, 393. Plumbago, 183. Plumbum, 357. Nigrum, 357. Polyhalite, 294. Polymerism, 206. Polymers, 206. Porcelain, 352. Portland cement, 325. Positive electrode, 121. Index. Potash, 297. Name, 294. Red prussiate, 387. Yellow prussiate, 387. Potassium, 293. Alum, 349. Antimonyl tartrate, 274. Bichromate, 366. Bromide, 230. Carbonate, 297. Chlorate, u, 12, 296, 297. Chloride, 295. Chloroplatinate, 394. Chromate, 366. Chromium sulphate, 368. Cyanide, 198, 298, 315. Dichromate, 366. Discovery, 284. Ferricyanide, 387. Ferrocyanide, 198, 387. Hydroxide, 297, 298. Hypochlorite, 139. Iodide, 232, 233. Manganate, 371. Name, 294. Nitrate, 155, 295. Nitrite, 295. Permanganate, 370. Preparation, 294. Preservation, 294. Properties, 294. Relation to life, 298. Salts and starch, 298. Salts at Stassfurt, 293. Silicate, 258. Sulphate, 298. Sulphocyanate, 198, Tartrate, 418. Test, 294. Pottery, 352. Powder, gun, 296. Smokeless, 428. Precipitate, 45. Precipitation, 44. Prefix, centi-, 437. Prefix, continued. Deca-, 437. Deci-, 437. Hecto-, 437. Hydro-, 91, 95. Hypo-, 91. Kilo-, 437. Milli-, 437. Per-, 91, 95. Press cake, 296. Pressure, normal, 18, 19. Priestley, n, 16, 18, 55, 64, 140, 158, 161, 445. Primary phosphates, 269. Print, photographic, 313. Problems, 21, 30, 49, 59, 86, 108, 132, 146, 165, 180, 201, 224, 234, 254, 264, 277, 283, 300, 318, 330, 342, 353, 364, 37 2 391, 395, 436, 439, 440. Based on equations, 107. Producer gas, 25. Products, 83. Addition, 204. Substitution, 203. Proof spirit, 410. Propane, 203, 409. Properties of matter, I, 2. Propyl, 409. Propylene, 202, 204. Proust, 77, 445. Prout, 398, 446. Prussian blue, 388. Prussiate of potash, red, 387. Yellow, 198, 387. Prussic acid, 198. Puddling, 379. Pulp, paper, 429. Purification, water, 39. Purifiers, gas, 212. Purple of Cassius, 317. Putty, 323. Pyrite, 386. Pyrogallic acid, 431. Pyroligneous acid, 415. 582 Index. Pyrolusite, 369. Pyromorphite, 357. Pyrophosphates, 269. Pyrophosphoric acid, 269. Pyrosulphuric acid, 252. Pyrrhotite, 373. Quadrivalent elements, 176. Qualitative analysis, 50, 242. Quantitative analysis, 50. Quantitative significance of equations, 104. Quantivalence, 176. Quartation, 315. Quartz, 255, 256. Quartzite, 256. Quicklime, 324. Quicksilver, 338. Quinine, 433. Quinquivalent elements, 176. Radical, 89, 150, 198. Organic, 406. Valence, 177. Rain water, 37. Ramsay, 68, 69, 446. Rational formula, 407. Rayleigh and Ramsay, 68. Reaction, 3. Acid, 90. Alkaline, 92. Chemical, 83. Illustrating equation, 106. Neutral, 94. Realgar, 272, 273. Red fire, 329. Hematite, 374. Paint, 340, 361, 384. Lead, 360. Liquor, 350, 417. Reduction, 15, 28, 55, 357. Process for lead, 358. Reducing agent, 28. Flame, 222. Reference books, 450. Refining petroleum, 208. Relative humidity, 66. Respiration, 16, 191. Retorts, coal, 210. Reverberatory furnace, 281, 282. Reversion, 271. Rhigolene, 208. Rhodium, 392. Rhodocroisite, 369. Rinmann's green, 390. River water, 38. Rochelle powder, 290. Rock, crystal, 255. Phosphate, 271. Rocks, 258. Decayed, 265. Phosphorus from, 265. Silicates, 255. Roll sulphur, 238. Rosaniline, 431. Rosendale cement, 325. Rose's metal, 275. Rouge, 384. Royal water, 1 60. Rubidium, 284, 299, 413. Ruby, 347. Ore, 301. Rum, 411. Run, water gas, 213. Rusting of iron, 383. Rutherford, 63, 64, 446. Ruthenium, 392. Saccharose, 423. Safety lamp, 221. Sal ammoniac, 151. Saleratus, 290. Salicylic acid, 431. Sal soda, 289. Salt, 94. Acid, 96. As glaze, 352. Basic, 96. Cake, 288. Common, 286, Index. 583 Salt, continued. From White Sea, 287. Glauber's, 292. Microcosmic, 269. Preparation of common, 287. Springs, 228. Saltpeter, Chili, 231, 292. Source, 295. Salts, 94. . Action on litmus, 94. Ammonium, 150. And ionization, 129. Epsom, 333. Ethereal, 419. Formation, 94. General properties, 88. Haloid, 225. In ocean, 38, 287. Nomenclature, 95. Normal, 96. Organic, 419. Smelling, 152. Sand, 255, 256. And hydrofluoric acid, 228. Blast, 257. Sandstone, 256. Saponification, 422. Sapphire, 347. Satin spar, 326. Saturated compounds, 177. Hydrocarbons, 203. Point of air, 66. Solution, 44. Scandium, 401. Scheele, 16, 18, 64, 133, 265, 446. Scheele's green, 273. Scrubber, 212. Seal, 210. Sea water, salts in, 287. Silver in, 308. Secondary phosphates, 269. Seidlitz powders, 290, 418. Selenite, 326. Selenium, 252. Series, homologous, 202. Paraffin, 203. Serpentine, 331. Shell, in limestone, 322. Rock, 322. Shot, 360. Sicily, sulphur from, 236. Siderite, 373, 387. Siemens-Martin process for steel, 382. Silica, 255. And plants, 257. Deposition, 258. From springs, 258. Hydrated, 256. Soluble, 258. Silicates, 257, 258. Siliceous sinter, 258. Silicic acid, 257. Silicides, 258. Carbon, 117. Silicified wood, 256, 257. Silicon, 255. Bronze, 305. Carbide, 117. Tetrafluoride, 228, 257. Silicon dioxide, 255. Properties, 256. Varieties, 255. Silver, 308. Acetate, molecular weight, 1 70. Alloys, 308, 311. Amalgam, 309. Amalgamation process, 309. Bearing lead, 308. Brick, 310. Bromide, 312. Chloride, 308, 309, 312. Coins, 311. Compounds, 312. Compounds and light, 312, 313. Determination of atomic weight, 171. Distribution, 309. German, 306. Glance, 308. 5 8 4 Index. Silver, continued. Halogens, solubility, 252. History, 308. Horn, 133, 308. In sea water, 308. Iodide, 312. Metallurgy, 309, 310. Name, 308. Nitrate, 312. Ores, 308. Oxidized, 311. Plating, 311, 312. Production, 308. Properties, 310. Pure, 310. Separation from gold, 315. Specific heat, 173. Sterling, 311. Sulphides, 308, 311. Tarnishing, 311. Test, 312. Water, 338. World's supply, 308. Silverware, blackening, 242, 311, Simplest formula, 104, 175. Sinter, siliceous, 258. Sirius, 23, Sirup, table, 426. Slag, 281, 324, 375. Phosphate, 271. Slaked lime, 324. Slate, 343. Smalt, 390. Smelling salts, 152. Smelting, 281. See Metallurgy. Smithsonite, 334. Smokeless gunpowder, 428. Snow crystals, 35. Soap, 420, 422. And hard water, 327. Boiling process, 423. Cold process, 423. Hard, 422. Soft, 422. White, 422. Soap, continued. Yellow, 423. Soapstone, 331. Soda, 289, 290. Ash, 289. Baking, 290. Cooking, 290. Crystals, 289. Washing, 289. Water, 42, 90, 193. Sodium, 284. Acetate, 417. Alum, 349. Aluminate, 348, 349. Amalgam, 292, 339. And water, 24, 51. Arsenate, 273. Arsenite, 273. Bicarbonate, 195, 289. Carbonate, 284, 288, 289. Chloride, 286, 287. Cyanide, 286, 293. Dioxide, 293. Discovery, 284. Hydroxide, 290, 291, 292. Hypochlorite, 139. Hyposulphite, 138, 252. lodate, 230. Iodide, 319. Lactate, 418. Manganate, 372. Manufacture, 284, 285. Monoxide, 293. Name, 284. Nitrate, 292, 293. Oxides, 286. Peroxide, 286, 293. Preservation, 286. Properties, 285. Silicate, 258. Stannate, 357. Sulphate, 292. Sulphide, 288. Sulphite, 243. Test, 141, 286. Index. 585 Sodium, continued. Thiosulphate, 252. Tungstate, 369. Uses, 286. Soft coal, 185, 189. Water, 37, 327. Solder, 356, 360. Soldering, 263. Solid carbon dioxide, 193. Solids, solution, 43. Table, 44. Soluble glass, 257. Silica, 258. Sulphate, test, 251. Solute, 41. Solution, 41, 126. And chemical action, 47. And electrolysis, 126. Boiling point, 127. Freezing point, 127, 128. Gases, 41. Labarraque's, 139. Liquids, 43. Nature, 48. Saturated, 44. Solids, 43. Supersaturated, 45. Terms, 41. Thermal phenomena, 47. Solvay process for sodium carbonate, 289. Solvent, 41. Universal, 43. Souring, 138. Sour milk in cooking, 418. Specific gravity of metals, 279. Specific heat, 172. Law, 172. Table, 173. Spectra, 402. Nebulae, 404. Stars, 404. Spectroscope, 23, 402. Discovery by, 284, 404. Spectrum, 401. Spectrum, continued. Absorptive, 403. Analysis, 401, 403. Banded, 402. Bright line, 402. Dark line, 402. Sunlight, 403. Speculum metal, 306. Spelter, 335. Sperrylite, 392. Sphalerite, 334. Spiegel iron, 369, 378. Spinel, ruby, 347. Spinels, 347. Spirit of salt, 140. Spirits, hartshorn, 147. Spongy platinum, 392, 393. Springs, mineral, 37, 42. Stable refuse, 271. Stack, 210. Stahl, 1 6, 446. Stalactite, 321. Stalagmite, 321. Stamp, mill, 280. Standard conditions, 19. Wax candle, 216. Stannic chloride, 357. Oxide, 356. Stannous chloride, 356. Stannum, 354. Starch, 426, 427. And potassium salts, 298. Test, 232, 427. Stas, 171, 398, 446. Stassfurt deposits, 133, 228, 261, 293, 331- . Steam, 36. Stearic acid, 417. Stearin, 420. Candles, 422. Steel, and coke, 190. Bessemer, 381. Chrome, 366. Crucible, 381. Harveyized, 380. 586 Index. Steel, and coke, continued. Manufacture, 380. Nickel, 389. Open hearth, 383. Properties, 380. Tempering, 380. Uses, 383. Sterling silver, 311. Stibine, 274. Stibium, 274. Stibnite, 274. Still, 40, 379. Stone, artificial, 258. Ice, 350. Stoneware, 352. Stove polish, 183. Strass, 260. Stream tin, 356. Striking back, Bunsen flame, 220. Strontia, 328. Strontium, 328. Carbonate, 328. Hydroxide, 328. Nitrate, 328. Oxide, 328. Sulphate, 328. Sulphide, 329. Test, 329. Structural formulas, 178. Stucco, 327. Sublimate, 151. Corrosive, 340. Sublimation, 151, 440. Subnitrate of bismuth, 276. Substitution, 3, 203. Products, 203. Sucrose, 423. Suffix, -ate, 95. -ic, 91, 144. -ide, 95. -ite, 95. -ous, 91, 144. Sugar, 423. Barley, 424. Beet, 424. Sugar, continued. Brown, 424. Cane, 423, 424. Fermentation, 410. Fruit, 425. . Granulated, 425. Grape, 425, 426. Kinds, 423. Of lead, 363. Of milk, 425. Raw, 424. Refining, 425. Term, 423. Test, 426. White, 424. Suint, 293. Potassium carbonate from, 297. Sulphates, 235, 251. Acid, 251. Important, 251. Normal, 251. Test, 141, 251. Sulphides, 238, 241. Color, 242. Native, 235. Solubility, 242. Sulphites, 245. Acid calcium, 245. Acid sodium, 245. Sodium, 243. Sulphur dioxide from, 243. Sulphur, 235. Action with heat, 238. Allotropic modifications, 239. Amorphous, 239, 240. And metals, 238. And silver, 311. Burning, 245. Compounds, 240. Crystallized, 239. Dioxide, 242, 244, 245. Extraction, 236. Flowers, 238. Formation, 235. Forms, 239. Index. 58? Sulphur, continued. Free, 235. In human body, 236. In United States, 236. In volcanic districts, 235. Kiln, 236. Milk of, 240. Monoclinic, 239. Native, 235. Orthorhombic, 239. Properties, 238. Purification, 237. Roll, 238. Source, 236. Springs, 37, 235. Trioxide, 245, 246. Use, 240, 252. Vapor density, 238. Water, 37. Sulphuretted hydrogen, 240. Sulphuric acid, 246. And organic matter, 250. And water, 250. Chemical changes in making, 248. Concentration, 249. From pyrites, 386. Fuming, 251. Impurities, 363. Manufacture, 246, 248, 249. Nordhausen, 252. Plant, 247-248. Properties, 250. Reduction, 250. Test, 251. Uses, 251. Sulphuric ether, 414. Sulphurous acid, 244, 245. Anhydride, 245. Sulphocyanic acid, 198. Sun, elements in, 23, 404. Sunlight and carbon dioxide, 194. Chemical action, m. Nitric acid, 156. Superheater, 213. Superphosphate of lime, 271. Supersaturated solution, 45. Supporter of combustion, 15. Sylvite, 294. Symbols, 81. And atomic weights, 103. Chemical, 8. Latin, 8. Table, 448, 449. Synthesis, 3, 50. Table salt, 287. Tables, atomic weights, 448, 449. Borax bead colors, 262. Composition of coal, 186. Composition of natural waters, 38. Equivalents, 100. Famous chemists, 447. Important elements, 6. lonization, 127. Latin symbols, 8. Metric equivalents, 438. Metric system, 437. Metric transformations, 438. Multiple proportions, 78, 79. Periodic, 399. Solubility of carbon dioxide, 42. Solubility of solids, 44. Specific heats, 173. Uncommon elements, 7. Water in food, 32. Talc, 331. Tallow, 421. Tannic acid, 432. Tannin, 432. Tanning, 433. Tar, 213. Extractor, 212. Well, 210. Tartar, crude, 418. Emetic, 274, 419. Tartaric acid, 418. Tea, 433. Tellurides, 314. Tellurium, 252. Compounds, 314. 588 Index. Temperature and luminosity, 218. Kindling, 113, 218. Low, 204. Standard, 19. Tempering, 380. Temporary hardness, 327. Tension of water vapor, 36. Terms, electrochemical, 120. Terra cotta, 352. Tests, acetic acid, 419. Alcohol, 419. Aluminium, 347. Antimony, 275. Arsenic, 273. Barium, 329. Bismuth, 276. Borax bead, 262. Boron, 261. Cadmium, 337. Calcium, 328. Carbon, 189. Carbon dioxide, 192, 325. Chloride, 144. Chromium, 367, 368. Cobalt, 390. Copper, 304. Gold!, 317. Hydrochloric acid, 144. Hydrogen sulphide, 242. Ions, 129. Iron, 388. Lead, 363. Lithium, 298. Manganese, 372. Marsh's, for arsenic, 273. Nickel, 389. Nitrates, 159. Nitric acid, 159. Potassium, 294. Silver, 312. Sodium, 286. Starch, 427. Strontium, 329. Sugar, 426. Sulphate, insoluble, 251. Tests, continued. Sulphate, soluble, 251. Sulphuric acid, 251. Zinc, 337. Tetrads, 176. Tetragonal crystals, 441. Theine, 433. Theory, 75. Atomic, 79. Electrolysis, 125. Electrolytic dissociation, 125, 126. Thermal equation, 112. Thermometers, 439. Thiosulphate, sodium, 252. Thomas-Gilchrist process for steel, 382. Tiles, 352. Tin, 354. Alloys, 356. Amalgam, 339, 356. Block, 355. Crystals, 356. Dioxide, 354, 356. Foil, 356. History, 354. Interaction with metals, 355. Metallurgy, 354. Ore, 354. Oxymuriate, 357. Plate, 355. Production, 354, 356. Properties, 355. Purification, 355. Stone, 354. Stream, 356. Uses, 355. Tinkel, 261. Tinware, 355. Tobacco, 433. Toluene, 202, 430. Toluidine, 431. Toning, in photography, 313. Topaz, 347. Travertine, 322. Triacid base, 94. Index. 589 Triads, 176. Tribasic acid, 92. Triclinic crystals, 442. Trivalent elements, 176. Tungsten, 369. Turnbull's blue, 388. Turquoise, 347. Tuscany, boric acid from, 261. Tuyeres, 376. Type metal, 360. Water, 89. Univalent elements, 176. Unsaturated compounds, 177. Hydrocarbons, 204, 206. Uranium, 369. Salts, 369. Specific heat, 173. Urea, 405. Valence, 176. Classification by, 397. Representation, 407. Valentine, Basil, 246. Van Helmont, 196, 446. Van't Hoff, 446. Vapor density, 169. And molecular weight, 168. Iodine, 232. Mercury, 339. Sulphur, 238. Zinc, 336. Vapor tension, 36. Varec, 231. Vaseline, 209. Vegetable matter and coal, 184-185. Vein mining, 315. Venetian red, 384. Verdigris, 417. Vermilion, 340. Vinegar, 90. Preparation, 415. Quick process, 416. Wood, 415. Vital force, 405. Vitriol, blue, 307. Green, 385. Oil of, 92, 246. White, 337. Volatile alkali, 93, 149. Volta, 119. Voltaic cell, 119. Volume equation, 175. Volumetric, 53. Composition of air, 64. Composition of water, 53, 55, 57. Washing soda, 289. Washington monument, cap, 345. Water, 31. Analysis, 39. And chlorine, 51. And hydrogen, 50. And oxygen, 51. And sodium, 24, 51. As solvent, 32, 33. Baryta, 329. Boiling point, 36, 439. Chalybeate, 37, 387. Chlorine, 135. Composition, 25, 27. Density, 34. Distilled, 40. Drinking, 39. Electrolysis, 52, 123. Expansion, 34. Freezing, 34, 439. From burning hydrogen, 27. Function in nature, 32. Gas, 25, 196, 213, 214, 215. Glass, 25^. Gravimetric composition, 55, 57. Hard, 37, 327. Hardness, 327. Hydrogen sulphide, 241. Industrial application, 33. In food, 31, 32. In human body, 32. In liquid state, 31. In vegetables, 31, 32. 590 Index. Water, continued. Javelle's, 139. Lithia, 298. Mineral, 37. Natural, 37. Occurrence in nature, 31. Ocean, 38. Of crystallization, 45, 46. Physical properties of pure, 33. Purification, 39, 371. Quantitative composition, 53. Rain, 37. River, 38. Silver, 338. Soda, 42. Soft. 37, 327. Type, 89. Underground, 37. Volumetric composition, 53, 55, 57. Water vapor, 31, 36. Condensed, 31, 36. In atmosphere, 62, 66. Watt, 55. Wax, paraffin, 209. Welding iron, 379. Weldon, mud, 370. Process, 134, 370. Welsbach light, 222. W T et process, 47, 282. Whetstone, 256. Whisky, 411. White arsenic, 272. Cast iron, 378. Lead, 361. Magnesia, 370. Metal, 306. Paint, 242, 336, 362. Vitriol, 337. Whitewash, 326. Whiting, 323. Willemite, 334. Willson, 1 1 6. Winds, 62. Wine, in, 297. Witherite, 329. Wohler, 343, 405, 446. Wood alcohol, 409. Ashes, 297. Charcoal, 187. Petrified, 256, 257, 258. Preserving, 337. Silicified, 256, 257. Spirit, 409. Vinegar, 415. Wood's metal, 275, 337. Worm, condenser, 40. Wrought iron, 378. Xenon, 69, 404. Yeast, 410. In bread-making, 427. Yellow paint, 367. Zinc, 334. Alloys, 306, 336. Blende, 334. Carbonate, 334. Chloride, 122, 337. Deposits, 334. Determination of atomic weight 173- Dust, 335, 336. Hydroxide, 337. Metallurgy, 334. Ores, 334. Oxide, 334, 335, 336, 362. Production, 334. Properties, 335. Silicate, 334. Smelting, 334. Sulphate, 336. Sulphide, 334, 336. Test, 337. Uses, 336. Vapor density, 336. White, 336. Zincates, 335, 337. Zincite, 334. Zero, absolute, 439. THIS BOOK IS DUE ON THE LAST DATE STAMPED BELOW AN INITIAL FINE OF 25 CENTS WILL BE ASSESSED FOR FAILURE TO RETURN THIS BOOK ON THE DATE DUE. THE PENALTY WILL INCREASE TO SO CENTS ON THE FOURTH DAY AND TO $1.OO ON THE SEVENTH DAY OVERDUE. 15 f93 vTieio^ APR 25 1940 9 1940 WUM* ^^ :,^ - ;! 1 LD 21-95m-7,'37 YB 16941 THE UNIVERSITY OF CALIFORNIA LIBRARY