IRLF SB 35 373 Edmund O'Neill ELEMENTS OF GENERAL CHEMISTRY WITH EXPERIMENTS. BY JOHN H. LONG, M. S., Sc. D., PROFESSOR OF CHEMISTRY AND DIRECTOR OF THE CHEMICAL LABORATORIES IN THE SCHOOLS OF MEDICINE AND PHARMACY OF NORTHWESTERN UNIVERSITY. CHICAGO : E. H. COLEGROVE, SALES AGENT. 1898. Entered according to act of Congress, in the year 1898, BY JOHN H. LONG, in the office of the Librarian of Congress, at Washington. INMEMORIAM, PREFACE. In the preparation of this book the author has used a part of his work, " Experimental and Analytical Chemis- try," now out of print, and has added enough new matter to make of it a complete text-book of elementary general chemistry, sufficient for the wants of college students beginning the subject. In too many instances the student is introduced to qualitative analysis as his first laboratory work, and this is followed by gravimetric analysis to complete a course. This plan certainly gives the beginner a distorted idea of the relative importance of analytical chemistry in the study of the science; for the beginner a knowledge of the proper- ties of a substance, the methods of its preparation and its uses is far more important than acquaintance with methods of separation, and general illustrative experiments should, therefore, be made the foundation work in the laboratory. It is the belief of the author that much that is demon- strated by the teacher in the classroom may profitably be repeated by the student in the laboratory. Repetition is necessary to fix elementary principles thoroughly in the mind of the beginner. The list of experiments here offered embraces the work of this character which has been required in the author's classes during the past ten years. Most of these exercises are simple and easily performed; others are longer or more complex, and are therefore described in considerable detail, but all of them may be performed by the aid of comparatively simple apparatus, 889793 and all, it is believed, illustrate important facts or prin- ciples. In the descriptive part of the book the author has kept in mind the fact that it is intended for beginners, few of whom expect to become specialists in chemistry, and he has, therefore, made the presentation of matter as practical as possible. Some important substances and technical processes are described more fully than is usually thought necessary in an elementary book. No chemical theory is introduced in the earlier chapters, but after the student has been made familiar with important principles by experi- ment it is gradually presented. In explaining the atomic theory an attempt has been made to show in a very ele- mentary manner the important steps, historically, in its development. It is believed that this method will give the student the clearest insight into a subject which is, at best, hard to grasp and which is seldom mastered. The author wishes to acknowledge the very valuable assistance rendered him by his friend and colleague, Dr. Charles H. Miller, in reading proofs and in other ways helping in the publication of the book. THE AUTHOR. CHICAGO, 1898. TABLE OF CONTENTS. CHAPTER I. Introductory I 1 CHAPTER II. Oxygen, Hydrogen and their Compounds 30 CHAPTER III. Chlorine and Hydrochloric Acid. Theoretical Considerations 66 CHAPTER IV. Compounds of Chlorine with Oxygen. Bromine, Iodine, Fluorine and their Compounds.. 90 CHAPTER V. Nitrogen and the Atmosphere. Gas Problems 106 CHAPTER VI. Compounds of Nitrogen 120 CHAPTER VII. Sulphur and its Compounds, Selenium and Tellurium 146 CHAPTER VIII. Silicon and Boron and their Com- pounds 172 CHAPTER IX. Phosphorus and Arsenic and their Com- pounds 183 CHAPTER X. Carbon and some of its important Com- pounds 202 CHAPTER XL Atomic and Molecular Weights 231 CHAPTER XII. Classification of the Elements. Gen- eral Properties of the Metals and their Salts 249 CHAPTER XIII. The Alkali Metals: Lithium, Sodium, Potassium, Rubidium and Caesium. Ammonium Compounds 271 CHAPTER XIV. The Copper Group: Copper, Silver and Gold 288 CHAPTER XV. The Alkali Earth Group: Beryllium, Magnesium, Calcium, Strontium and Barium. The Spectroscope 310 CHAPTER XVI. Zinc, Cadmium and Mercury 328 CHAPTER XVII. Boron, Aluminum, Gallium, Indium, Thallium, Scandium, Yttrium, Lanthanum and Ytterbium 330 CHAPTER XVIII. The Carbon Group: Carbon, Silicon, Germanium, Tin and Lead. The Titanium Group: Titanium, Zirconium, Cerium and Thorium 344 CHAPTER XIX. The Nitrogen Group: Nitrogen, Phos- phorus, Vanadium, Arsenic, Columbium, Antimony, Tantalum and Bismuth 354 CHAPTER XX. The Chromium Group: Chromium, Molybdenum, Tungsten and Uranium. Relations to the Oxygen Group 362 CHAPTER XXI. Manganese and its Relations to the Halogen Group 369 CHAPTER XXII. The Iron Group: Iron, Nickel and Cobalt 375 CHAPTER XXIII. The Platinum Group: Ruthenium, Rhodium, Palladium, Osmium, Indium and Plati- num.. ..389 CHAPTER I. INTRODUCTORY. IN BEGINNING the study of chemistry in the labora- tory or classroom the student should learn to consider each experiment performed as a; question , and th& result obtained its answer. Chemistry is^fe^mitieritlyran^exper- imental science in which matte* under A^t^in, conditions is the subject of investigation. By experiment and observation we seek to determine the properties of this matter, to divide it into groups, to analyze and decide what is simple and what compound, to find the action which one kind of matter exerts upon another and how each one behaves under the influence of heat, light, electricity and other forces. We seek also to find the simplest and best means of producing different kinds of matter, and to discover tests by which they may be always recognized. This knowledge, with more to be acquired, when prop- erly classified and arranged in a consistent system, consti- tutes the science of chemistry. The beginner can best obtain acquaintance with this science by his own experi- ments in the laboratory under the guidance of an instruct- or. Much can be and must be learned from books, it is true, but the knowledge which is most satisfactory and most lasting when acquired is that which the student gathers by direct contact with the thing under study. In many lines observation alone brings but limited knowledge. For instance, of the air or of the water every- where around us, we would know indeed but little if un- aided by experiment. When we 'make an experiment on an object we take the thing, in a sense, within our grasp and look at it from different sides, placing it under new and varied conditions, and by so doing learn many of its 2 GENERAL CHEMISTRY. important qualities and peculiarities. Asking ourselves how it would behave under certain conditions, we make the experiment and find out. In chemistry we study mat- ter as undergoing change. We are acquainted with matter in three general forms or conditions, the gaseous, liquid and solid states, and we shall first give our attention to a brief consideration of these. The Three States of Matter. Many kinds of matter are found to exist in the three forms mentioned,, but k>r each substance there is a condi- tion in : which' it is^mos-t stable and most usually found. Th charge: from; one condition to the others is generally most r^raiiiit^'broti^htalDOutby a change of temperature; a low temperature being favorable to the maintenance of the solid condition while a high temperature aids in the forma- tion of gases or vapors. We have in water a familiar illustration of a substance well known in the three condi- tions, but many other common substances can readily be made to pass from one of these conditions to the others, as can be shown by experiment. Ex. I. Let the student apply heat to a test-tube one-third filled with sulphur. At a temperature of about 115 C. it melts to a yellow liquid which grows darker by application of more heat and becomes viscid. At a still higher temperature the viscid mass becomes thinner, and finally boils at a temperature of about 450 C. Application is made of this fact in the refining of sulphur by distillation. In this experiment the vapor of the sulphur usually ignites at the mouth of the test-tube and burns with a pale blue flame, forming sul- phurous oxide, as will be explained later. Ex. 2. In a somewhat narrow test-tube melt two or three grams of camphor. This passes from the solid to the liquid condition at a temperature of 175 and boils at 204. Vapors are given off even at low temperatures, from which it follows that in experimenting with quite small pieces of camphor the middle or liquid condition may be over- looked If the tube taken is long enough, 10 to 12 Cm., the vapor from the boiling liquid will condense on the upper and cooler part. With iodine and several other bodies the phenomena of vaporization are very similar. Iodine melts at about 115, but gives off vapors at a lower temperature. So rapid is GENERAL CHEMISTRY. 3 vaporization above the melting point, that the temperature of actual ebullition cannot be accurately observed. It is above 200. With ammonium chloride, or sal ammoniac, the behav- ior is different. The substance gives off no vapors at the ordinary temperature, but readily at high temperatures. Ex. 3. In a test-tube heat some of the ammonium chloride in fine crystals. Observe that it does not melt, but at a sufficiently high tem- perature gives off dense white vapors, which soon condense to crystal- line grains. We have here the passage from the first to the third state without liquefaction. Many of our common and best known substances can- not be obtained in the state of vapor, and some not even in the liquid condition, because they surfer decomposition when strongly heated. The red oxide of mercury is a good illustration of this, as it separates into mercury and oxygen by heat. Common limestone breaks up when heated, yielding quicklime and carbonic acid gas. Common salt and potassium chlorate may be liquefied, but suffer decomposi- tion when heated to higher temperatures. The so-called organic substances are those in which passage through the three states can be most readily ob- served, as the temperature of vaporization is in general much lower here than among the inorganic compounds. The change of all bodies from the liquid condition to that of a gas or vapor depends not only on temperature, but also on the pressure on it, that of the atmosphere usually. Variations of the atmospheric pressure cause a change in the temperature to which a substance must be brought to change it from a liquid to a vapor. Vaporization follows at a lower temperature by de- crease of the air pressure on the heated liquid. Hence it is that many substances which cannot be distilled under the ordinary atmospheric pressure without decomposition can be easily and safely distilled in a partial vacuum. Solutions. When common salt is thrown in water and the mixture stirred the salt gradually disappears, leaving finally a clear 4 GENERAL CHEMISTRY. liquid which in appearance cannot be distinguished from the water. Many other substances behave in the same manner, for instance, sugar, saltpeter, soda, borax and sal ammoniac. We apply the name solution to the clear mix- tures of these substances with water. The sugar, salt and other substances are said to be soluble in water. They are soluble, also, in other liquids. In such cases the particles of the solid seem to distribute themselves among the liquid particles, and in every instance there is a limit to the power or capacity of the water for dissolving the solid. In the illustrations given, as in manyothers, the solvent and body dissolved exert no decomposing action on each other, because the two can be readily separated and obtained in their original conditions. To make this point plain let the student make the following experiment : Ex. 4. Into some distilled water stir clean, pure salt, a little at a time, until the water becomes saturated^ that is, until it will no longer take up any more of the salt. In 50 cubic centimeters at the ordinary temperature we can dissolve in this manner about 18 grams of salt. Next pour about half of the solution into an evaporating dish, place this on wire gauze or on a sand-bath (sand in an iron dish) over the low gas flame from a Bunsen burner and heat slowly. The water gradually disappears or passes off in the form of vapor, leaving at last the dis- solved salt as a clear white crystalline mass. Many substances not soluble in water are soluble in alcohol, ether, chloroform or other liquid, and usually without change; that is, by evaporation of the menstruum the substance may be recovered as was the salt in the above experiment. But other bodies dissolve only by decomposition. For example, marble is not soluble in water, but it can be quickly dissolved by action of certain acids. Ex. 5. In a small beaker take a few grams of chalk or powdered marble (commonly called marble dust). Add water and stir or shake the mixture thoroughly. Then allow to settle, and, as far as can be de- termined by the eye, it will be noticed that the marble remains undis- solved. Next add, a few drops at a time, some hydrochloric acid and the escape of gas which follows shows that an important change is taking place. Gradually add more acid until the effervescence, after shaking, ceases. There should be left now a clear or nearly clear liquid or solution, and by evaporating this in a small dish or beaker, the slight excess of acid employed in making it will be driven rff. What remains is soluble in water, while the original marble was not. GENERAL CHEMISTRY. 5 The action of the acid here has been to convert an in- soluble body into one readily soluble in water. We can- not properly speak of the solution as a solution of marble, as this substance is no longer present. As another simple illustration of solution effected by conversion into a new substance, the action of acids on many metals may be referred to. It will be shown later that iron, zinc and other common metals dissolve readily in hydrochloric or sulphuric acid. During the action of the acid on the metal a gas escapes and there is left dissolved a combination, termed a salt, of a part of the acid with the metal. In a mixture holding a body in solution and something insoluble in suspension, the latter may be separated by filtration, that is, by passing the liquid through an appara- tus termed a filter, which holds the particles not actually in solution. As a filtering medium, paper, sand, porous stone, felt, and other substances may be used. In illus- tration of this, make the following experiment: Ex. 6. In a beaker mix some common salt and clean marble dust. Pour on water and shake thoroughly. Allow to subside and then pour the liquid on a paper filter. (For method of making a paper filter the instructor must be consulted.) To the residue in the beaker add more water, stir again and pour through the same filter. Finally, wash the residue itself from the beaker onto the filter and allow it to drain. When dry it will be recognized as the original marble dust. The liquid which passed through the filter, or filtrate, on evaporation yields the salt. In this experiment the fine pores or openings in the paper permit the passage of the liquid and the salt dis- solved in it, but not the passage of even the finest parti- cles of the undissolved marble dust. Filtration is one of the most common operations of analytical chemistry. Bodies differ, when soluble, very greatly in the extent of their solubility. While common salt will dissolve in less than three times its weight of water, at the o r dmary temperature, or cane sugar in about half its weight of water, there are required for gypsum nearly 400 and for morphine nearly 1,000 parts of water. Indeed, it may be said that no substance is absolutely insoluble in water, but where the degree of solubility is so small that several 6 GENERAL CHEMISTRY. thousand parts of water are required for solution it is cus- tomary to speak of the body in question as practically insoluble. In many classes of investigations, however, even very slight degrees of solubility must be taken into consideration. The temperature of a liquid has in most cases a marked influence on its solvent power. It usually happens that the solubility of a substance is increased by increase of temperature, but this is not always the case. We know of a few common substances which are actually less soluble in warm liquids than they are in cold. A striking illustra- tion of the change of solubility with change of tempera- ture is shown by the behavior of niter or saltpeter with water. At a temperature of 20 C, that is at a common room temperature, 100 parts by weight of water dissolve about 31 parts of the niter, but at a temperature of 100 C., that is at the temperature of boiling water, nearly 250 parts may be dissolved. The following experiment may be made by way of illustration: Ex. 7. In a test-tube take about 10 Cc. of water at the laboratory temperature. Add to it some powdered saltpeter, a little at a time, close with the thumb and shake after each addition until the water be- comes saturated, that is until it will no longer dissolve the added salt- peter. It will be observed that as the solid goes into solution the temperature of the liquid becomes lower. Now gradually heat the solution in the Bunsen burner flame and from time to time add more of the saltpeter. In this way it will be seen that a clear solution may be made containing many times the weight of the substance dissolved in the cold. When it has become saturated at the boiling heat set it aside to cool slowly. After a time the test-tube will be found to contain a mass of crystals deposited by the cooling of the liquid. Other interesting examples of the same effect of tem- perature may be referred to. At 0, 100 Gm. of water dis- solves about 4 Gm. of crystallized potassium alum, but at 100 the same weight of water dissolves over 350 Gm. of the alum. At 0, 100 Gm. of water dissolves about 5 Gm. of crystallized oxalic acid, but at 100 nearly 350 Gm. may be dissolved. A solution is said to be saturated at a certain temper- ature when it contains as much of a substance as it will hold at this temperature. It is a peculiarity of many substances, GENERAL CHEMISTRY. 7 however, that they may be temporarily dissolved in water or other liquid in amount greater than can be permanently held at the same temperature. A solution so produced is said to be supersaturated. This condition is most readily attained by dissolving a salt by the aid of heat until the solution becomes saturated at a high temperature. On carefully pouring off some of the clear hot liquid it may often be cooled to the air temperature and kept a long time without precipitation. Let, however, a small crystal of the dissolved substance be dropped into the cool liquid, a precipitate from the solution will appear and settle out until the amount remaining dissolved is just sufficient to constitute a normally saturated solution at the given low temperature. The change from the state of supersatura- tion to that of normal saturation is here brought about by addition of some of the same salt that is dissolved, and in general, supersaturation in a solution at a given tempera- ture may be detected in this manner. It may be detected also by other means. The phenomenon is one of so much importance that it will be illustrated by experiment. Ex. 8. In each of three perfectly clean beakers holding about 250 Cc. dissolve 50 Gm. of pure, powdered, crystallized sodium sulphate in 25 Cc. of distilled water. Heat to 30-35 to complete the solution. When a clear solution is obtained cover each one of the beakers with a piece of paper and set them aside to cool in a perfectly still place. When the three solutions are quite cool test them as follows: Remove the papers and into one beaker drop a small crystal of the sodium sul- phate. By means of a glass rod rub the bottom and sides of the second beaker, while the contents of the third maybe poured out into a dry beaker. In each case the equilibrium of the solution is destroyed and a precipitate of the dissolved salt settles out. The substances which most readily form supersaturated solutions are those which combine with large amounts of "water of crystallization." This term will be explained later. The alums and borax resemble sodium sulphate in this behavior. Crystallization. In the foregoing the formation of solutions has been explained and it has been shown that the dissolved sub- stance may often be readily separated or recovered from 8 GENERAL CHEMISTRY. the solution. In the case of bodies not decomposed by the menstruum this is most readily effected by evaporation, as illustrated by the recovery of salt or sugar dissolved in water. The widest application of this fact is made in the arts in the production or purification of numerous impor- tant substances. Concentration of a solution is often suf- ficient to throw out the dissolved substance. The solid very frequently assumes what is termed the crystalline form as it leaves the solution, and this is gener- ally the case when it is deposited slowly, as by the gradual cooling of a liquid. In the experiment with saltpeter the formation of crystals was shown, but the phenomenon can be better illustrated by the use of another substance. Ex. 9. Dissolve 25 Gm. of powdered alum in 75 Cc. of water by aid of heat. Filter the hot solution into a clean beaker which may then be set aside in a quiet place for spontaneous evaporation. Several hours or over night should be given for this and at the end of the time large crystals of alum will be found. To prevent the too rapid cooling of the solution immediately after its preparation the vessel containing it may be wrapped in cotton, or better still, in felt. Slow cooling favors the production of large, well-formed crystals. Ex. 10. Prepare strong solutions of copper sulphate, or blue vitriol, and chrome alum by dissolving about 15 Gm. of each in a small quantity of warm water. For each about 50 Cc. of water should be used. Filter the solutions into small clean beakers, which set aside in a quiet place protected from dust for several days. The dissolved salts will begin after a time to crystallize out as the solvent water evaporates spontaneously. Slow evaporation and a low temperature favor here, as before, the formation of large crystals. Very perfect crystals of many substances may be made by a slight modification of the above experiment. If in very strong solutions of the alums, blue vitriol, potassium dichromate or potassium ferrocyanide, for instance, a small crystal of the same substance be suspended by means of a fine thread this crystal will serve as a nucleus around which a deposit forms as the solutions become concen- trated by spontaneous evaporation. A very pretty effect is obtained by growing in this manner a good crystal of chrome alum. This is then suspended in a cold satu- rated and clear solution of common potash alum, when the growth continues, the potash alum being deposited on the GENERAL CHEMISTRY. 9 chrome alum. Many substances which are isomorphous, that is, have the same crystalline form, can be crystallized together in this manner. The process of crystallization is employed in many in- dustries and in chemical investigation on the small scale for the purification of substances. This can be illustrated by an experiment. Ex. ii. Dissolve some crude common salt in hot water to make a saturated solution. Filter this hot into an evaporating dish and con- centrate a little by heat. On allowing now to cool, some of the salt will separate in pure white form. By repeating the operation on the liquid remaining, the mother liquor, it is called, further crops of crystals may be obtained. As salt is found in nature its natural contaminations are usually substances much more soluble than it is. These are therefore left in the mother liquor. The first crops of crystals are the purest, but if the concentration be carried too far, the salt obtained may be mixed with these impurities. The filtration at the beginning of the operation above was intended to remove insoluble sub- stances only. By fractional crystallization it is often possible to sepa- rate two or more substances from a solution. This is true where the substances dissolved differ greatly in their de- grees of solubility. Sodium chloride may be separated from sodium nitrate in this manner, and copper sulphate from potassium sulphate. In concentrating solutions of either one of these mixtures the least soluble substance will begin to separate first. The first fraction may be very nearly pure. By continuing the concentration and crystal- lization from the mother liquors the last fractions obtained may be nearly pure crystals of the most soluble substance. By dissolving now, the first fraction obtained in water and crystallizing again the crop of crystals obtained will be nearly or quite pure in some cases. The mother liquor is used as the solvent for the second fraction which yields now a fresh portion of the least soluble salt and holds more of the most soluble. By a continuation of this method the most soluble constituent can be concentrated in a solution practically free from the others, and then crystallized itself. 10 GENERAL CHEMISTRY. Water of Crystallization. Many substances in crystallizing from aqueous solution unite with a part of the water, holding it in the form de- scribed as water of crystallization. Other substances crys- tallize in the anhydrous form that is, they hold no water. Common salt and saltpeter are familiar illustrations of bodies belonging to the second group, while blue vitriol, alum and Glauber's salt are common substances which contain water of crystallization. Blue vitriol is a combina- tion of copper sulphate with water, and when the sub- stance is powdered and strongly heated the water is driven off, leaving the copper sulphate in the anhydrous or pure form. This pure copper sulphate is no longer blue, but white. In ordinary usage the term copper sulphate is understood to refer to the common blue crystallized com- pound. The behavior of this substance when heated may be shown by experiment. Ex. 12. In a narrow test-tube heat a few grams of powdered blue vitriol in the gas flame, but not to a high temperature. To avoid too great heating the tube may be moved backward and forward, and turned meanwhile between the fingers, at a point some distance above the hottest part of the flame. Four-fifths of the water held by the sub- stance is given off at a temperature not far from 100 C. f and may be oeen as vapor in the tube. The sulphate is left as a bluish white pow- der after the escape of the vapor. The heat may now be increased so as to drive off the remaining water. With care this can be done without breaking the tube, when it will be seen that the residue is a nearly pure white powder. If the test-tube be now allowed to cool and some water added, the powder will immediately unite with a part of it, becoming blue again. On exposure to the air, this white powder takes up mois- ture enough to give it a blue color in a very short time. Some common substances contain so much combined water that when heated they appear to melt and assume the liquid form. Ordinary potassium alum and sodium thiosulphate (commonly called hyposulphite of soda) show this phenomenon. Ex. 13. In a test-tube carefully heat some small crystals of the sodium thiosulphate. It will be seen that they melt very readily and at a low temperature. If strongly heated, water is driven off and can be recognized. The liquid obtained by melting the salt in its water of crys- tallization behaves as a supersaturated solution, which can be shown as GENERAL CHEMISTRY. 11 follows : After liquefying the substance close the tube with cotton or a cork, and stand it in a quiet place where it can cool down without any jar or agitation. Under these conditions the substance remains as a liquid. If now the stopper be removed and a minute crystal be dropped into the tube, the contents solidify immediately. Some substances holding water of crystallization can be dehydrated without decomposition. We have illustra- tions of this in blue vitriol, borax, crystal soda, alum, Glauber's salt and others. From this it appears that the hydrated and anhydrous forms of these substances are equally stable. But other substances holding water are stable only in this condition and decompose when an attempt is made to separate their water. Attention will be called later to the exact chemical composition of sub- stances crystallizing with water, which can be best illus- trated by means of formulas. Precipitation. It has been shown that many solid substances can be dissolved or brought into solution by means of water or other liquid. The converse of this will now be illustrated. That is, it will be shown that solutions can often be made to give up their dissolved substance by other methods than by evaporation or crystallization. This is commonly accomplished by the process termed precipitation. By this we understand a process in which a dissolved solid or some part of it is rendered insoluble and settles out from, or is precipitated from the solution. Precipitates are usu- ally heavier than the liquid from which separated and therefore settle to the bottom of the containing vessel. A substance in solution may be rendered insoluble in several ways, for instance, by change of temperature, by adding to the solution a second liquid in which the dissolved sub- stance is insoluble, or, most commonly, by converting the dissolved substance into a new one, insoluble in the men- struum, by addition of some decomposing reagent. As an illustration of precipitation by change of temper- ature the following test may be made : Ex. 14. To about a gram of calcium tartrate in a test-tube add 10 12 GENERAL CHEMISTRY. cubic centimeters of caustic soda solution. The solid dissolves by shak- ing. When a clear solution is obtained boil it and observe that a gela- tinous precipitate forms. This consists of the calcium tartrate, ren- dered insoluble by increase of temperature. We have numerous illustrations of precipitation by ad- dition of a second liquid to the solution, and the methods are frequently applied in practical analysis. The follow- ing experiments will serve as illustrations: Ex. 15. Dissolve some crystallized ferrous sulphate (common green vitriol) in warm water and filter the solution to make it perfectly clear. To some of this clear liquid in a test-tube add an equal volume of alcohol and observe that a precipitate of the sulphate in small crystals settles out. A more satisfactory result can be obtained by pouring the sulphate solution into the alcohol, shaking thoroughly and then setting the mixture aside several hours. The ferrous sulphate is readily soluble in water, but not in alcohol; hence on adding the latter to the solution the iron compound settles out unchanged. Many salts may be precipitated from their aqueous solutions by alcohol in the same manner. Ex. 16. By aid of heat dissolve dextrin or other gum in water. To the solution add some alcohol and observe the precipitation of the gum. Allow this to subside, which may require hours; pour off the liquid as far as possible and add pure water. This brings the gum into solution again, indicating that the addition of alcohol had rendered it only temporarily insoluble. Ex. 17. Dissolve common rosin or colophony in alcohol. Pour some of the clear solution into a test-tube and add an equal volume of water. Precipitation of the resin substance follows. Gums are soluble in water usually, but not in alcohol. Resins and many similar bodies are soluble in alcohol but not in water. Hence precipitation takes place in one case by adding alcohol to the aqueous solution and in the other by adding water to the alcohol solution. In the illustrations given the substance precipitated separates from its solution in practically unchanged con- dition. Precipitation here is not accompanied by decompo- sition. In the great majority of cases, however, what is termed precipitation depends on change of chemical composition, and is brought about by adding to a solu- GENERAL CHEMISTRY. 13 tion something, usually a solution of another substance, which is capable of producing a new and insoluble body with that already present. The new body formed must therefore settle out as a precipitate. The nature of this change can be made plain best by a few simple experi- ments. Ex. 18. To some dilute solution of blue vitriol (copper sulphate) in a test-tube add an equal volume of solution of ammonium sulphide and shake. From the mixture of the blue copper solution and the nearly colorless or yellow sulphide solution we obtain a black substance which is evidently not the original copper sulphate. This belief is con- firmed by filtering the contents of the test-tube. A yellow or brownish liquid passes through the paper while a bulky black precipitate remains. By pouring water on this precipitate it fails to dissolve, showing its marked difference from the vitriol. This black substance is known as copper sulphide, and in many important properties is quite unlike the copper sulphate from which it was produced. Ex. 19. To a solution of alum in a test-tube add some ammonia water and shake the mixture. A very bulky gelatinous precipitate forms and gradually settles to the bottom of the test-tube. It can be separated by filtration from the liquid in which it was suspended, and when mixed with pure water is found to be insoluble. In appearance and in characteristic properties this substance is very different from the origi- nal alum. It is called aluminum hydroxide. Ex. 20. By the aid of heat dissolve a few grams of powdered chalk in weak, hydrochloric acid contained in a test-tube or small beaker. When solution is complete boil a few minutes, and filter if the liquid is not perfectly clear. We have now a solution containing, not the chalk, because the acid decomposed that, but calcium chloride, a new substance. That we have here a substance distinct from the chalk can be shown by evaporating some of the solution to dryness in a small porcelain dish, heated on a sand bath. The appearance of the residue, and the fact that it dissolves in water while the chalk does not, show the distinct nature of the body. Now, to the remainder of the solution not evaporated add a little ammonia water, enough to make it impart a blue color to red litmus paper after stirring, and then some solution of am- monium carbonate. This will produce a fine white precipitate which settles readily to the bottom of the vessel. After it has stood some hours, pour off the liquid as far as possible, and collect what remains on a filter. Allow the fine white precipitate to drain thoroughly and then pour water over it. When this has run through add water a second time and wait for this to drain. Then stand the filter aside and allow the white precipitate to become thoroughly dry, which may require sev- eral days. On examination of what remains it will be found to have the appearance and properties of the original chalk. In dissolving the chalk in the acid it was observed that a gas was given off and this gas is 14 GENERAL CHEMISTRY. known as carbon dioxide or carbonic acid gas. One element of the chalk, however, certainly remained behind, because a solid substance was found on evaporating the solution. It appears, therefore, that in the formation of chalk again the ammonium carbonate solution must have restored in precipitation just what was lost when the chalk went into solution. Pure chalk is known chemically as calcium carbonate. An important peculiarity of precipitation in general is shown by these examples. We have first a body in solu- tion with its particles uniformly distributed among those of the solvent. A condition of equilibrium exists of such a nature that any tendency of the particles of a heavy body to sink or of a light body to rise and float is exactly overcome. Bodies in solution resemble gases in this respect, that their particles tend to separate and fill all available space uniformly. We have seen that this condition of equilibrium maybe destroyed in several ways by change in temperature, by addition of a new liquid in which the dissolved body is in- soluble, or by addition of a certain solution. This solution must contain a substance capable of forming a third sub- stance insoluble in the mixed solutions. We may have an aqueous solution of a substance, A, and a second aqueous solution of a substance, B, but it does not follow that the product of the action of A on B should also be soluble in water. It often happens that the product of A and B is insoluble. For example, sodium sulphate and barium chloride are both easily soluble in water, but on mixing their solutions we obtain one of the least soluble of known substances, barium sulphate. The formation of a precipitate in a liquid is not an in- stantaneous operation, although in some cases the interval between the addition of the precipitant and the formation of a precipitate is very short. The precipitation of barium sulphate is an illustration. But more time is required for the completion of many other reactions, as will be seen by the following experiment : Ex. 21. Let the student pour some dilute solution of magnesium sulphate into each of three tes't-tubes. (This solution may be made by dissolving 5 Gm. of the crystallized substance in 100 Cc. of water.) To the first test-tube add solution of barium chloride; a precipitate forms immediately, apparently. To the second add an equal volume of a GENERAL CHEMISTRY. 15 solution of calcium chloride containing 3 Gm. in 100 Cc. A precipitate will slowly form. To the third test-tube add an equal volume of a 10 per cent solution of ammonium chloride, then some ammonia water, and finally a few drops of solution of sodium phosphate. In time a crystalline precipitate will appear. . The formation of this precipitate may be aided by rubbing the sides of the test tube with a glass rod, and the insoluble substance settling out appears first in the form of minute glistening specks, which grow larger and finally disclose a crystalline structure. It is evident from this that the phenomenon of precipi- tation is a complex one. The substance we recognize as a precipitate is not immediately formed, but is a growth, the particles we see being formed by the aggregation, proba- bly, of an almost infinite number of smaller particles. The building up or development of these larger particles is often greatly aided by application of heat. In the precip- itation of barium sulphate in the above experiment the precipitate remains for a long time suspended in the mixed liquid. By having both liquids warm it settles sooner, while if the mixture be boiled after precipitation the white precipitate will settle very rapidly, giving evidence of the heaviness and compact form of its particles. A loose pre- cipitate of barium sulphate, as it is produced in cold solu- tions, cannot be easily filtered. The particles appear to be so fine that they can pass through the pores of ordi- nary filter paper. After thorough boiling, however, filtra- tion is generally easy, the particles becoming coarse enough to be retained on the filter. Ex. 22. As a further instructive illustration of slow precipitation the following experiment may be made. In a test-tube mix 5 Cc. of a cold solution containing about 6 Gm. of tartar emetic in 100 Cc. with an equal volume of a dilute sodium carbonate solution containing about 1 Gm. in 100 Cc. Apply heat to the mixture and observe that a white precipitate forms immediately. Now repeat the experiment using the same solutions in the same quantities, but have both as cold as possible before mixing and pour the soda solution into the other very slowly, and with little agitation. Close the test-tube with a cork and leave it in a quiet place; under these conditions hours may elapse before the slight- est trace of precipitation appears. On shaking the tube, pouring out the contents or slightly warming, a precipitate begins to form and soon becomes heavy. The behavior here recalls that already observed in the experiments with supersaturated solutions, and the mixed 16 GENERAL CHEMISTRY. liquid just before precipitation was in a supersaturated condition; the subsidence of the precipitate relieves this. In the great majority of cases of precipitation the time in which supersaturation can be said to exist is extremely short, so as to escape observation. Precipitates are distinguished from each other by color, size of particles, apparent density of particles, rapidity of formation and subsidence, degree of insolubility and in many other ways. No two precipitates are exactly alike and we have therefore in the phenomenon of precipitation something of value for the recognition of substances, In analytical chemistry precipitation plays a very important part as a means of separation and identification. In chem- ical industry many substances are secured by precipitation from solutions containing them. In the pages to follow these and other applications will be abundantly illustrated. Distillation. It was explained in the beginning of this chapter that many substances are capable of existing in three forms, as solids, liquids and gases, or vapors. The conversion of a solid or liquid into a vapor is usually termed vaporization and may take place spontaneously, or, commonly, by the application of heat. The operations of vaporization and subsequent conden- sation of the vapor to the liquid or solid condition again, taken together, constitute what is termed distillation. When sufficient heat is applied to water in a flask, it boils and steam is formed which escapes from the mouth of the flask. The production and escape of the steam alone do not con- stitute distillation, but if the neck of the flask is closed with a perforated cork, or rubber stopper, through which a long glass tube, bent downward after leaving the flask, passes, some or perhaps nearly all of the steam will con- dense to form water which may be collected from the end of the tube. The flask and bent glass tube constitute a rude distil- ling apparatus which can be readily constructed by the student and used for the following experiment. See Fig. 1. GENERAL CHEMISTRY. 17 Ex. 23. Into the neck of a glass flask holding about 300 Cc. fit a good cork or rubber stopper having a perforation at least three-eighths of an inch in diameter. Next select a piece of glass tubing just wide enough to fit the hole in the stopper snugly and about three feet long, and melt the rough ends in the flame of the Bunsen burner to remove the sharp edges. Then about four inches from one end of the tube make a bend by heating it in a broad flame until it is soft enough to be bent so that the shorter limb makes an angle of about 60 with the longer. (The method of working glass tubing must be learned from the instructor.) This shorter limb passes through the perforation in the stopper. Pour into the flask about 150 Cc. of water, add some salt, enough to give a strong taste, and then a little indigo solution or other highly colored liquid. Insert the stopper with its bent glass tube, sup- port the flask on a sand-bath or wire gauze by means of iron rings or FIG. 1. clamps and then apply heat slowly, below the sand or gauze. The water in the flask becomes hot and finally begins to boil. Steam passes up into the bent tube and then condenses readily, if the heat applied is not too strong. Allow a few drops of the condensed liquid to fall from the end of the tube and then collect what follows in several perfectly clean test-tubes. It will be observed that the colored liquid in the flask yields a colorless distillate, and also that the latter is free from salty taste. To the water collected in one of the tubes add a few drops of solution of silver nitrate. No change should follow. To some water containing salt, as poured into the flask, add silver nitrate and observe that a heavy white curdy pre- cipitate forms. These experiments show that the salt which gives the characteristic taste and the white precipitate with the silver solution does not pass over with the steam. It is not readily volatile. The sub- stance of the colored liquid is likewise nonvolatile. 18 GENERAL CHEMISTRY This experiment illustrates the manner of separation of a volatile from a nonvolatile substance in general. Water and other liquids are commonly purified by distillation; that is, they are in this manner separated from solid sub- stances they hold in solution. Practically, the very simple distillation apparatus used in the experiment cannot often be employed. In most cases the condensation of the vapor would be quite imperfect. Instead of the simple glass tube a more elaborate condenser is usually attached to the flask or still, and the forms best known should be FIG. 2. on exhibition in the laboratory. In the great majority of cases in practice condensation is effected by passing the vapor through a straight or worm tube surrounded by flowing cold water. Forms of distilling apparatus are shown in the figures 2 and 3. In practical laboratory work the operation of distillation is a very common one. By it liquids may often be separated from solids, volatile solids from nonvolatile, easily volatile liquids from such as are not readily vapor- ized, or solids volatile at a low temperature from those volatile at a high temperature. In these latter operations GENERAL CHEMISTRY. 19 the method known as fractional distillation is often em- ployed. The principal applications of fractional distilla- tion are in organic chemistry, but a simple illustration may be given here. An approximate separation of water, alcohol and ether may readily be made, because these substances boil at very different temperatures. Ether boils at 35 C., alcohol at 78.5 C. and water at 100 C. Therefore if a mixture of these substances be distilled from a flask and the distillate collected in small portions or fractions it is evident that FIG. 3. the first fractions will consist mainly of ether and the last of nearly pure water, while in the fractions collected near 80 C. the alcohol will be in excess. A sharp separation is not possible, because the ether, beginning to boil at 35, carries with it in the form of vapor some alcohol and even a little water. In turn the vapor of alcohol carries with it some water vapor, so that the fractions are far from pure at first. A familiar illustration of the application of fractional distillation on the large scale is found in the refining of 20 GENERAL CHEMISTRY. crude petroleur.i, which consists of a mixture of many liquids of different boiling points. A full explanation of the phenomenon of fractional distillation would be out of place at this time, but may be found in the larger works on organic chemistry, a subject for later study. Chemical and Physical Changes. In our experiments on the precipitation of barium sul- phate and several other substances we had an illustration of what is termed a chemical change. In the melting of iodine or sulphur we had an illustration of a physical change. The distinction between chemical and physical changes will be made plain by a few simple experiments. Ex. 24. Heat some pieces of bright copper or iron wire in the hot flame of the Bunsen burner. Observe that the surfaces become tar- nished and that by repeating the operation several times a dark, brittle scale is formed which can be easily rubbed or scraped off with a knife. Next heat a small piece of magnesium wire in the Bunsen flame. When it becomes quite hot it burns with a white, dazzling light, giving off a white, cloud-like substance which finally settles down as a powder. Finally, heat a piece of platinum wire in the same hot flame. It will be seen that while held above the lamp the metal becomes very hot and bright red, but no evidence of scaling or formation of fumes is seen. On removing the wire from the heat it resumes its former color, and as far as can be seen it is in no manner different from what it was before heating. These simple experiments are very instructive. By heating the copper or iron it was evident that something new was formed with properties different from those of the original metal. The black scale scraped from the iron or copper is brittle and hard, while the metals are ductile. The white powder formed from the magnesium is evi- dently quite distinct from the metal and it becomes appar- ent that in the operation of heating something has been lost by the metals or absorbed by them which changes them into new substances. In an early period of chemical study it was held that under the influence of heat metals lost something. It is now known that instead of losing weight the copper, the iron and the magnesium take up something from the air which converts them into new sub- GENERAL CHEMISTRY. 21 stances, with an increase instead of a loss in weight. This absorption of something from the air with increase in weight constitutes a radical change in the substance under experimentation, a change in which its characteristic prop- erties disappear, giving place to equally marked properties in the new substance. The identity of magnesium is so completely lost in the white powder formed by burning that the recognition of the relation of the two substances is regarded as one of the triumphs of early chemical investigation. Changes as far reaching as these, changes involving frequently a loss of identity, are spoken of as cJiemical changes. Even superficial examination shows that no radical al- teration takes place in the platinum during the heating operation. The change there was merely a temporary one, involving no real loss of identity. Such changes are termed physical. Ex. 25. In a test-tube mix some flowers of sulphur with fine cop- per turnings. Gradually apply heat to the mixture. At first the sul- phur melts and becomes very dark colored. As the temperature grows higher a point is reached where a combination suddenly takes place between the sulphur and the copper which is shown by the glowing of the latter. The copper seems to burn in the atmosphere of sulphur in the tube. After this experiment the tube is allowed to cool and may be broken. In place of the bright, ductile copper, a black, brittle body is found, which evidently has but few of the properties of the original metal. The substance formed here is termed copper sulphide, and in its production we have a typical chemical change. Ex. 26. Pour some solution of blue vitriol into a beaker, and add to it a little dilute sulphuric acid. Next add a few small fragments of granulated zinc, and allow the beaker to stand half an hour. It will soon be recognized that a change is going on in the beaker, as the zinc becomes coated with a dark, spongy mass, in color suggesting the cop- per. It will be observed, also, that the blue color of the liquid gradu- ally becomes fainter, and finally that it may disappear entirely. (This depends on the amount of zinc taken.) By removing the spongy mass from the beaker, washing and drying it, the properties of copper maybe recognized. Meanwhile it should be observed that the zinc has wholly or in part disappeared We have here a very curious chemical change, which will find a fuller explanation later. But it may be said now that the zinc appears to go into the solution, while the 22 GENERAL CHEMISTRY copper of the blue vitriol or copper sulphate solution is precipitated. At any rate, the identity of the solution is destroyed with the loss qf its copper. Ex. 27. Over some small nails or tacks in a beaker pour some dilute sulphuric acid. Very soon an evolution of gas is observed, and after a time the metal will have disappeared. A light green liquid results, and this evidently contains the iron in a dissolved form. We have, in fact, a solution of green vitriol or ferrous sulphate, which could be separated by crystallization. It will not be necessary to multiply instances, as enough has been given to show what is characteristic in so- called chemical changes. Iron, under the action of a high heat or by treatment with an acid, becomes changed mate- rially by conversion into something else which is not iron, but which contains iron. Under the action of a strong magnet iron becomes likewise changed, but, as we know, only temporarily. On the removal of the magnet the iron assumes its original nature and important properties^ and gives little or no evidence of the physical change through which it has passed. A careful study of the common chemical changes or reactions shows that we can make three general divisions of them. We have first, reactions of decomposition in which a single substance is broken up or decomposed so as to yield two or more other substances. We have many illus- trations of this. For example, in "burning" lime the common rock known as limestone is strongly heated until it becomes decomposed, yielding a residue called quicklime which on addition of water becomes slaked lime, and a gas called carbonic acid gas or carbon dioxide. This is generally allowed to escape. The action of the heat here is to effect disintegration, but it adds nothing in the form of matter to the limestone. A simple experiment will be given in which the decomposition of a substance is easily shown. Ex. 28. In a small test-tube heat a few grams of r^d mercuric oxide to a high temperature The substance darkens and finally begins to break up, as may be readily shown by two phenomena. In the test- tube above the heated powder a deposit of fine metallic globules col- lects and this is easily recognized as mercury itself. If while a strong GENERAL CHEMISTRY. 23 heat is being applied a glowing splinter be held just within the mouth of the test-tube it will burst into flame and burn with great brilliancy. This shows that in addition to the metallic globules furnished by the red compound a gas is liberated, for only a gas could exhibit the behav- ior just mentioned. This experiment will be taken up again. We have here a characteristic reaction of decomposi- tion without the aid of outside matter in which we obtain from a heavy red powder a silvery liquid, mercury, and a gas, oxygen. Such reactions are frequently termed analyt- ical reactions because they consist in an analysis or break- ing up of something. We have next reactions of just the opposite character, that is, reactions in which two or more substances com- bine to form a third body. Our experiment on heating the copper and sulphur is an illustration of reactions in this group. At a high temperature the two substances were shown to combine, forming a new compound called copper sulphide or sulphide of copper. A second illustra- tion is furnished by the rusting of iron. Here the metal combines with something from the air (oxygen), produc- ing oxide of iron. It should also be mentioned that under certain conditions the reactions described above by which limestone and the red oxide of mercury were each sepa- rated into two substances are reversible. That is, the lime and carbonic acid gas may be combined to form lime- stone, and metallic mercury and oxygen to form the red oxide. Many such reactions are known and they are some- times called synthetical reactions. The most important and numerous of our chemical changes belong to a third group, however. Here two or more substances react on each other to produce two or more new substances. Several illustrations of such de- compositions were given above under the head of precipi- tation. Another may be given here. Ex. 29. Take a few grams of granulated zinc in a beaker and pour over it some dilute sulphuric acid. An effervescence begins immedi- ately showing the escape of .a gas, in some manner produced by the action of the zinc on the acid. This is an evidence of the formation of at least one new substance, because the gas can be neither the zinc nor the acid. It will be readily seen that the acid dissolves the zinc, that is, that a solution is formed, and when the action is complete, which is 24 GENERAL CHEMISTRY. shown by the disappearance of the metal, pour some of the solution into a small porcelain evaporating dish on a sand-bath, and apply heat to drive off everything volatile. Finish the concentration in a fume closet, applying finally a strong heat. A white residue will be left which is plainly neither zinc nor sulphuric acid. We have, therefore, in this case the formation of a gas and a white solid substance from a metal and an acid liquid, which can readily be shown to be wholly volatile. The gas is hydrogen and the white solid is zinc sulphate. Ex. 30. In a test-tube take about 10 Cc. of strong " sugar of lead" solution (solution of lead acetate). Heat to boiling and observe the odor. In another test-tube take an equal volume of dilute sulphuric acid, boil and observe the odor. Mix the hot liquids. A white precipi- tate forms which certainly does not resemble either one of the original substances. It will be noticed also that the mixed liquid emits a strong odor of vinegar or acetic acid. To show more clearly what has hap- pened allow the precipitate to settle in the test-tube and then pour the liquid above it through a filter. Heat the filtered liquid to the boiling point and observe that the odor is very strong and characteristic. Then add water to the residue in the test-tube, warm and pour the mixture on the same filter, and wash it several times by pouring on water. That this residue is not the sulphuric acid is evident, that it is not the lead acetate is shown by the fact that it is not soluble in the water poured over it, while the lead acetate is, readily. We have, therefore, in this case the production of an insoluble residue and a volatile liquid suggest- ing vinegar. The residue is known as lead sulphate and the volatile liquid is acetic acid. It can be readily shown that in many common precipita- tions two substances give rise to two new ones, but these illustrations are sufficient for the purpose at present. The fact that we obtained above two bodies from one by application of heat is sufficient proof of the compound nature of that body. It is plain that a body formed by the union of two must be compound, containing at least two component parts. Finally, when we obtain two new substances by the action on each other of two different bodies it is evident that one of them at least must be compound. These considerations can be illustrated by symbols as follows : Let AB represent a compound body which under cer- tain conditions is broken up into its component parts, A and B. Then we can write AB yields A -f B. GENERAL CHEMISTRY. 25 In the second case we have the reverse of this reaction, that is A + B yields AB. In the third case we go further and have evidently more component parts than A and B to consider. We have evidently C and D also, and we can express our result in a general way as follows : The bodies AB and CD act on each other and make AD and BC, or AB + CD yields AD + BC. At present no reason appears why we should not write instead of the above this: AB + CD yields AC -f BD. But later a meaning will be attached to these symbols which will render plain just what does take place. In cer- tain cases we can express our reaction in this manner: A + BC yields AB + C. In this instance only one of the bodies entering the reaction is considered as a compound one. This is BC, which the simple substance, A, decomposes into the new compound body, AB, and the new simple body, C. Chemists usually represent these changes by what are termed equations, as, A + B = AB AB = A -f B AB -h CD = AD + BC B + CD = D + CB. What is written to the left of the = sign represents that which is taken, and the result of the chemical change is shown on the right hand side of the sign. In these equations the letters A, B, C and D represent the elements or parts of compounds which take part in the reactions. Their full meaning will be explained later. Conditions of Chemical Change. The conditions under which chemical changes take place are different in different combinations. In some of 20 GENERAL CHEMISTRY. the illustrations given above it has been shown that cer- tain substances can be made to combine by the aid of heat, while in other cases, decomposition is effected by heat. These were cases, however, in which dry substances were taken for experiment. At the ordinary temperature such bodies enter into combination or decompose as a rule, but slowly. A few illustrations will be given in which solid bodies are combined by friction. Dry Reactions. The following three experiments are simple cases : Ex. 31. Rub together, in a mortar, about equal weights of corro- sive sublimate and potassium iodide. A bright red compound results, which is different from the substances giving rise to it, not only in color but in solubility, as may be shown. Add some water to the contents of the mortar and stir well. A red precipitate remains. This is a new compound, mercuric iodide. Ex. 32. Rub together in a mortar, minute quantities (a few milli- grams of each only} of sugar and potassium chlorate. The substances react on each other violently, producing an explosion. If large quanti- ties were used the experiment would be very dangerous. The chemical change taking place here results in the formation of bodies very differ- ent from the sugar or the potassium chlorate. Ex. 33. Mix in a mortar, by means of a piece of paper, or card- board, about equal weights of dry slaked lime and ammonium chloride. At first no change should be noticed, but on applying some pressure in mixing, as when the two substances are ground together with a pestle, a change rapidly takes place in which ammonia is liberated, as shown by the smell. The mass becomes moist as the rubbing is continued. The nature of this reaction will be explained later. Reactions in the dry way, as illustrated above, are in- teresting but not very common. A few have practical importance, but by far the greater number of chemical changes with which we are acquainted, take place in solu- tion. Reactions in Solution. Our experiments on precipi- tation are illustrations of these, but others may be given. Ex. 34. Dissolve very small amounts, as in Ex. 31, of mercuric chloride (corrosive sublimate) and potassium iodide in water and mix the solutions. The deep red precipitate results immediately. GENERAL CHEMISTRY. 27 Ex. 35. Mix the slaked lime and ammonium chloride mentioned in Ex. 33" with water, and warm gently. The strong ammonia odor soon appears. Vary this experiment then by using instead of the lime, solutions of caustic soda and sodium carbonate, which likewise liberate the ammonia. In these experiments the ammonium chloride is com- pletely decomposed, the volatile ammonia escaping. Ex. 36. Mix together on a piece of dry paper some sodium bicar- bonate (" baking soda ") and some dry powdered tartaric acid. As long as the mixture is kept perfectly dry no apparent change takes place. The substances do not seem to react on each other, and in fact the mix- ture may be kept dry almost indefinitely. But if it is thrown into a beaker and water added a lively effervescence begins, due to the escape of carbonic acid gas from the decomposed bicarbonate. The addition of water brings both substances taken into solution, in which condition they act readily on each other. The above is a typical experiment, as many changes take place in the same manner. The action of the common baking powders depends on the behavior here illustrated. A large number of substances seem to have no action on each other when mixed in the perfectly dry condition, but when dissolved mutual decomposition begins. In the above experiment it is not merely the soda which is altered, as shown by the escape of gas, but the tartaric acid suffers a change also. It is converted into a neutral body, that is, one without acid properties. It seems to be true that in solution the particles of dis- solved substance are brought into a condition in which tkey move with great freedom and may thus be brought into intimate contact with each other, which is not the case as long as they are in the dry form. In general, solutipn is favorable to chemical change, and we therefore, as far as possible, dissolve the substances we wish to combine with each other, for the production of new substances. Reactions of Gases. Not only have we reactions between solids and reactions between liquids, but we have also some well marked cases of reactions between gases. A few of these have practical importance. The following experiment will serve as an illustration: Ex. 37. By means of a glass rod place a drop of strong hydro- chloric acid on one side of the bottom of a dry beaker. Clean and dry 28 GENERAL CHEMISTRY. the rod, and with it put a drop of strong ammonia water on the opposite side of the beaker. The first drop contains hydrochloric acid gas, tne second ammonia gas. Some of each gas leaves the liquid in which it is dissolved and the two unite in the beaker, producing white fumes. These white fumes consist of ammonium chloride, a solid substance, which is finally deposited on the walls of the beaker. After placing the two drops in the beaker it should be covered with a glass plate. Several other gases combine readily in the same man- ner. In some cases the products formed are also gases; in other cases they are liquids, while sometimes, as in the above experiment, they are solids. In a few cases the combination takes place readily and spontaneously, but in other cases it must be brought about by special means. A mixture of hydrogen gas with oxygen gas may be kept at the ordinary temperature, but by application of heat and by other means the two gases combine with explosive violence, if in certain proportions, forming water. A mixture of hydrogen and chlorine gases may be kept in the dark, but" if brought into the sunlight a sudden com- bination or explosion follows, hydrochloric acid gas being formed. In later chapters other illustrations of gaseous combinations will be given. In all cases of chemical combinations, whether of sol- ids, liquids or gases, it has been found by experiment that the substances united combine in certain proportions only. For instance, it can be readily shown that 17 parts of am- monia gas, by weight, combine with exactly 36.5 parts by weight of hydrochloric acid gas. * If a larger amount of either one of these gases were taken with the given weight of the other, this excess would fail to go into union and would remain in the free state. In the reaction between hydrogen and oxygen 1 part by weight of the former combines with 8 parts of the latter. In combining hydrogen with chlorine, it is found that 1 part by weight of the former combines always with 35.5 parts of the latter, but not with more or less. In Ex. 27, it was shown that iron is dissolved by sulphuric acid. By careful attention to details it can be shown that the amount of iron which can be dissolved by a given weight of sulphuric acid is absolutely constant. GENERAL CHEMISTRY. 29 In Ex. 29 zinc is dissolved in the same acid, and proper tests show that the weight of the metal dissolved by a given weight of the acid is constant and greater than the weight of iron which can be dissolved in the same acid. From these illustrations it would appear, therefore, that in our chemical combinations we have quantitative as well as qualitative relations. It would be premature to attempt an explanation of these facts now. The student should bear them in mind and look for an explanation in later experiments. CHAPTER II. OXYGEN, HYDROGEN AND THEIR COHPOUNDS. OXYGEN. WE ARE now ready to begin the study of particular substances somewhat in detail, and will begin with the very important and common body known as oxygen. Occurrence. Oxygen is widely distributed throughout the animal, vegetable and mineral kingdoms, constituting about one-half of the total weight of everything we are acquainted with in and above the earth's crust. It makes up eight ninths of water by weight, and over one-fifth of the atmosphere. All the common rocks and clay contain it in combination, while in such common substances as sugar, starch, the fats, albumin and woody fiber it is an important constituent. History. The history of this remarkable body is inter- esting. While now easily recognized as a distinct sub- stance it must be remembered that the earlier chemists were without this knowledge. The curious properties which will be shown later to belong to oxygen were either overlooked or ascribed to something else. The atmosphere, which owes its most important properties to the oxygen present, was supposed to be a simple substance, and the common phenomena of combustion in air or oxygen were all wrongly interpreted. However, in 1774, Priestley, and, independently of him, Scheele, in 1775, isolated pure oxy- gen from compounds containing it and recognized it as the important element of the air. GENERAL CHEMISTRY. 31 In 1781 Cavendish announced the composition of water, and a little later the great French chemist, Lavoisier, gave the first rational explanation of the behavior of oxygen in combustion and respiration, and opened the way for the growth of the science of modern chemistry. Preparation. Although oxygen is abundantly present in well-known materials everywhere obtainable, we secure it in the pure state practipally from but few sources. It is obtained from the air at very slight cost by processes which cannot be explained at this point, but only where required in large quantities for certain manufacturing operations. When used for other purposes it is commonly made by the decomposition of certain compounds contain- ing it from water, from the red oxide of mercury, and from potassium chlorate, for instances. Two experiments will be here given to illustrate these operations: Ex. 38. Repeat Ex. 28 by heating a small amount of red mercuric oxide in a narrow test-tube. Observe that a strong heat is required for the decomposition of the substance, and that finally a glowing splinter held within the mouth of the test-tube rekindles and burns brightly. As already intimated, this phenomenon shows that a gas must be given off by the action of heat on the red compound. This gas is oxygen, and the simple experiment illustrates one of the first processes given for its preparation. This method o1 liberating oxygen is not a convenient one, and besides is very expensive. Experiment shows that 216 parts of the oxide of mercury yield only 16 parts of oxygen, and a high temperature is required to separate this from the mercury. The experiment has value only as an illustration, as by this method Priestley first secured the gas in pure condition. For laboratory uses we make oxygen generally by a process indicated by the following experiment: Ex. 39. In a test-tube heat about 10 Gm. of powdered dry potas- sium chlorate. Move the test-tube backward and forward through the flame, turning it meanwhile between the thumb and fingers so as to avoid cracking it. After a time the powder melts to a liquid, which by longer application of heat appears to boil. Gas bubbles are seen to escape from it, and if a glowing splinter is now held within the mouth 32 GENERAL CHEMISTRY. of the test-tube, it will soon burst into flame, as in the other case The high heat applied here decomposes the substance taken, and oxygen gas is one of the products. What remains in the tube will appear later. This process has certain drawbacks. A relatively high temperature is required for the breaking up of the chlorate and the reaction is somewhat slow. It may be modified, however, as follows: Ex. 40. Mix about equal weights, a few grams of each, of potas- sium chlorate and manganese dioxide. Heat the mixture in a test-tube and notice that the liberation of gas, which kindles the flame on a glow- ing coal, begins almost immediately. The substance does not melt, but at a relatively low temperature undergoes decomposition with separation of the oxygen. This decomposition is very rapid, as may be seen by the manner in which splinters of wood, once ignited, burn at the mouth of the test-tube. The last experiment shows us how oxygen may be prepared in quantity, which will now be tried. Ex. 41. Mix about 25 Gm. of dry powdered potassium chlorate with an equal weight of manganese dioxide, and transfer to a dry and perfectly clean glass flask, holding 200 to 250 Cc. By means of a per- forated stopper connect this flask with a bent delivery tube arranged as in the following figure. The further end of the delivery tube is bent upward slightly and dips beneath the surface of water contained in a pneumatic trough or earthenware bowl. The trough or bowl should be nearly filled with the water. The flask is supported on a sand-bath by means of a clamp or ring. On now applying heat by the burner, the mixture in the flask soon becomes hot and begins to decompose as shown by the escape of bubbles of gas from the end of the delivery tube. Method of Collecting the Gas. The gas is most readily collected by displacement of water, and the directions given for this operation here will answer for many later experiments. Have at hand several wide mouth bottles, holding about 250-400 Cc. each. Fill them quite full with water and then cover them with squares of glass in such a manner as to exclude all air bubbles. When this is done each bottle may be inverted by holding it with one hand, while the plate of glass is pressed down with the other. The mouth of the bottle in this position is brought beneath the surface of the water in the trough and the plate removed. The bottle remains full, its contents being held up by the atmospheric pressure, and this remains true in whatever position the bottle stands, provided its mouth is always beneath the surface of the trough water. It may therefore be held just over the end of the deliv- ery tube from the flask, and so catch the gas bubbles as they ascend. Instead of holding the collecting bottle in the hand it is much better to support it on a bridge of galvanized iron, which has a perforation in its center about 2 centimeters in diameter. The bubbles can pass through GENERAL CHEMISTRY. 33 this opening into the bottle above. When the bottle is full of gas move it to one side, its mouth still beneath the surface of the water, close with the square of glass and then lift it out of the water and stand on the table in an upright position. If the plate fits the bottle well, the gas will not soon escape. Now bring, in the same manner other bottles of water over the end of the delivery tube and collect enough gas for ah the experiments given below. The first bottle of gas collected may be contaminated by air from the generating flask, but the others should contain nearly pure oxygen. On completing the experiment the deliv- ery tube should be taken from the water before the lamp is removed. Why? FIG. 4. With the bottles of oxygen collected as just explained, the student is ready to make some simple experiments to illustrate important properties of the gas. Potassium chlorate furnishes 39.2 per cent ot its weight of oxygen when completely decomposed. At the ordinary temperature 1 Gm. yields about 300 cubic centimeters of gas. For the preparation of quantities of the gas a cop- per retort is commonly employed, and from this, on libera- tion, the gas is led into a large reservoir or holder. For certain purposes the gas requires some purification, unless made of absolutely pure materials. The purification can be easily effected by allowing the gas on leaving the gen- erating retort to bubble through a wash bottle containing 34 GENERAL CHEMISTRY. a solution of potassium hydroxide. The arrangement of generator, wash bottle and gas holder is shown in Fig. 5. One hundred grams of the chlorate mixed with an equal weight of manganese dioxide will furnish 30 liters of gas. A little of this should be wasted, however, to drive the air from the retort and wash bottle at the beginning of the operation. Before attempting to heat a large quantity of the mixed chlorate and black oxide, as explained above, a trial should FIG. 5. always be made with a small quantity in a test-tube. The black oxide has been occasionally found adulterated or accidentally mixed with charcoal or other form of carbon, and dangerous explosions have resulted from the careless use of such a product. The test-tube experiment would disclose such impurity if present. In the tests made in experiments 39 and 40 it was dis- covered that the gas supports combustion well, as shown by the rapid and brilliant burning of the splinter. A simi* lar phenomenon is shown in the next test. GENERAL CHEMISTRY. 35 Ex. 42. Attach a piece of charcoal to a bent wire, bring it to a glowing condition in the lamp flame, and plunge it quickly in a bottle of the gas. It burns brilliantly, throwing off showers of sparks. Pour some lime-water in the bottle at the end of the reaction, close with a glass plate and shake. A white precipitate is formed, owing to the com- bination of a constituent of the lime-water (calcium hydroxide), with the gas produced by the combustion of the charcoal. The gas is carbon dioxide. The precipitate consists of calcium carbonate. This experiment gives us a typical example of what is termed combustion. The oxygen in the bottle goes into intimate union or chemical combination with the hot char- coal, with production of a much higher temperature and ultimate destruction of the whole of the charcoal if the volume of oxygen is large enough. This union of the gas with the charcoal is attended by the escape of heat, and evidently intense heat, as indicated by the sparks. When carbon combines with pure oxygen gas, as in this case, the product formed is always carbon dioxide and this substance may be always recognized by the precipitate it yields with clear lime-water. It will be shown later that carbon diox- ide is formed by many other reactions. , Other bottles of the oxygen remain with which the fol- lowing experiments may be made : Ex. 43. Melt a small amount of sulphur in a deflagrating spoon (a small brass spoon with a long wire handle bent to make a right angle with the rim of the spoon), heat it until it begins to burn with a pale blue flame and then plunge it into a bottle of the gas, The sulphur burns here with a much brighter flame than in the air. In this combustion a gas is produced which has a very strong and characteristic odor. It is an oxide of sulphur, a combination of oxygen and sulphur, and is called, prop- erly, sulphurous oxide. A similar but weaker odor is noticed when some kinds of matches are burned. The bottle should be covered with a glass plate so that the contents may be saved for another experiment. Ex. 44. Carefully dry a "very smah piece of phosphorus between folds of filter paper, place it in a deflagrating spoon, ignite it by hold- ing for a second in the lamp, and thrust quickly into a bottle of oxygen. Immediate combustion takes place, an intense white light is produced, and the bottle soon fills with fumes of a substance known as phosphoric oxide. After a time these fumes settle to the bottom and sides of the 36 GENERAL CHEMISTRY bottle and mix with the moisture present. Cover the bottle with a glass plate and keep the contents for a second test. This experiment gives a very good illustration of the strong affinity , or liking, of the oxygen for phosphorus. The phenomenon is very similar to that presented in the burn- ing of charcoal, but the phosphorus-oxygen reaction is dis- tinguished by its greater intensity. The direction was given to heat the phosphorus to the burning point before putting it in the gas. This was done to save time rather than because of its necessity. If left to itself in dry condition in the air, spontaneous combustion takes place after a time. Because of this fact phosphorus must always be kept under water. In the experiments just given, substances which are readily combustible have been burned in the oxygen gas. Other bodies burn with greater difficulty, but in a manner none the less characteristic. This is true of many of the common metals, and as an illustration we will take iron. Ex. 45. Cut some fine soft iron wire into lengths of about 10 cen- timeters afid wrap a dozen or more of these into a bundle; at one end of this bundle spread the individual wires a little, and around the other end twist a piece of stronger wire to serve as a handle. Heat the loose end and dip it into some sulphur. Ignite this, and as soon as it burns well remove the cover from a jar of oxygen and hold the wire down into it. First the sulphur burns with greater brilliancy and finally the iron becomes hot enough to ignite too, burning in a manner reminding one of the charcoal combustion. The sulphur is used here merely to bring the iron up to the burning temperature. With coarse wire this experi- ment cannot be easily performed. In this combustion the product, like the others, is termed an oxide. After this reaction we have two substances in the bottle, but one of them is the sulphurous oxide produced by the burning of the sulphur. The iron oxide is left as a black scale or slag, which will not dissolve in water. Iron forms several other combinations with oxygen which will be referred to later. It remains now to make some tests with the products in the bottles after the burning of the carbon, sulphur and phosphorus. The oxides of carbon and sulphur are invisible gases; that of phosphorus appears for a time as a white cloud which subsides as mentioned. GENERAL CHEMISTRY. 87 Ex. 46. To the contents of each bottle add a little water, replace the cover and shake thoroughly. Whatever is present evidently com- bines with the water or goes into solution. In the water in each bottle dip a piece of blue litmus paper. The color changes to red. By means of a clean glass rod take up a drop of liquid from each bottle and touch it to the tongue. The drop from the carbon bottle imparts little or no taste, while those from the others are sour. We have here a common property of the bodies usually termed acids, as most of these have a sour taste and change blue litmus paper to red. The acid formed by the addition of water to the carbon bottle is evidently very weak. The simple experiments just detailed are exceedingly instructive. It can easily be shown that oxygen gas itself is free from sour taste and that it does not change the color of blue litmus paper. The addition of water does not alter this. The acid properties must therefore come by the union of the oxygen with the carbon, the sulphur and the phosphorus. Further, it can be shown that these ox- ides in the //restate are not acid bodies; they become such only after the addition of water. It would seem therefore that both oxygen and moisture are concerned in the pro- duction of acids, and we shall find later that in very many cases this is true. The name oxygen signifies acid producer, it being at one time supposed that all acids contained oxygen. The phenomena of combustion naturally attracted the attention of the ancient philosophers, but the explanations given of observed facts were very faulty. As intimated above, Lavoisier was the first to give a correct explanation of what appears to us now as very simple. His immediate predecessors in attempting to explain combustion and oxi- dation were led to the doctrine that metals, in burning, lost or gave up something, and this thing lost was called by them phlogiston. The fact that these metals grew heavier instead of lighter when burned was overlooked or consid- ered as of very little moment. The value of the balance in chemical investigation was yet to be shown. A few scien- tific men, recognizing, however, the importance of the change in weight, were forced to the absurd conclusion that phlogiston was endowed with a property of levity or nega- tive weight, so that when added to a body it made it weigh less instead of more. According to the phlogiston 38 GENERAL CHEMISTRY. view metals were mixtures of a calx or base with phlogis- ton, and application of high heat in the air liberated the light phlogiston, leaving the heavy calx. Lavoisier investigated all these points thoroughly and established the fact that when lead, iron, tin, mercury and other common metals are heated in the air they increa'se in weight by the absorption of oxygen ; and he was able to show in several cases that the gain in weight of the metal was equal to the loss in weight of the volume of air taken. He found, also, as shown above, that sulphur, phos- phorus and several other bodies burn in air and in oxygen producing gases, which in turn combine with water, form- ing acids. From his experiments he was led to believe that all acids must contain this gas, and hence he gave to it, in 1778, the name oxygen, or acid producer. Before this date it had been known by several fanciful names, as vital air, dephlogisticated air, and others. The term vital air indicates the importance of oxygen in the process of life. In an atmosphere devoid of oxygen animals die almost immediately, and experiment shows that a certain proportion of oxygen must always be present to preserve life. It has been mentioned that the atmos- phere in which we live contains about 20 per cent of oxygen. With a proportion notably below this, respira- tion in the higher animals becomes difficult; if it falls to a certain point, death must follow. More will be said on this question later. The oxygen taken into the lungs from the air serves for the combustion of carbonaceous substances in the tissues. Here, as in our simple experiments, carbon dioxide is formed, and this is thrown off from the lungs in the expired air. As the carbonaceous matters burned are very com- plex in structure, several other products are produced, some of which are excreted by the lungs, while others are thrown off by the kidneys or through the skin. In this combustion within the body, as in all others, heat is produced and the normal high temperature of animals is thus seen to depend on chemical combination or oxidation. Conditions of Oxidation. Some substances combine GENERAL CHEMISTRY. 39 with oxygen at the ordinary temperature, and if left exposed to the air burn or become quickly corroded. The presence of moisture in many cases greatly assists this spontaneous combination with oxygen, as is well illustrated by the rusting of iron in moist air. For other substances a certain elevated temperature must be reached before combination with oxygen will follow. Our experiments have shown that sulphur and charcoal burn readily in oxygen, but only after previous heating to what is called the kindling temperature. It is a well-known fact that at the ordinary temperature these two bodies are perfectly stable, showing no tendency to oxidize. When combustible bodies are heated to the kindling temperature in air or oxygen and burn, a certain definite amount of heat is always liberated. A gram of pure car- bon burning in a sufficient supply of oxygen liberates a constant number of units of heat, which term will be ex- plained later. The amount of heat liberated by the com- bustion of a gram of sulphur is much less than that formed from a gram of carbon, but is still a constant. The follow- ing short table gives in round numbers the units of heat liberated by combustion of 1 gram of the substances named, in pure oxygen. Hydrogen 34.000 units. Carbon 8,000 Sulphur 2,300 Phosphorus 5, 700 Zinc 1,300 Iron 1,570 Tin 1,230 Copper 600 Provided the end product is the same the amount of heat liberated is a constant whether the oxidation be rapid or slow, but the temperature reached by the burning body depends on the rapidity of the combustion. In slow oxi- dation, as in the rusting of a piece of iron, the elevation of temperature is not perceptible; but in rapid combustion heat enough may be liberated to melt the oxide formed. In the former case the heat liberated is dissipated by radi- ation or conduction, while in the second case the reaction 40 GENERAL CHEMISTRY. is so rapid that little time is given for loss in this way. Hence the elevation of temperature which follows. The subject of the heat of combustion is a very im- portant one in many scientific investigations, and also in the measurement of the calorific value of fuels. A knowl- edge of the heat of combustion of a number of simple sub- stances makes it possible to calculate the amounts of heat which will be liberated in the combustion of given weights of fuels of known composition. One liter of oxygen, measured at a temperature of 0C. and under a pressure of 760 Mm. of mercury (the so-called standard temperature and pressure), weighs 1.4298 Gm. Referred to dry air its specific gravity is 1.105G. It is but slightly soluble in water, 1 volume of the latter dis- solving 0.041 volume of oxygen at the standard temperature and pressure. Oxygen may be liquefied at a temperature of 118 at a pressure of 50 atmospheres. Uses of Pure Oxygen. At the present time oxygen gas as made by the chlorate process is employed for several purposes. It is made for inhalation to a slight extent, and for this purpose must be carefully purified. Much larger quantities are made for combustion with hydrogen or illu- minating gas in the production of the " calcium," 'Mime" or "Drummond" light. Oxygen gas is employed on a still larger scale in several manufacturing operations. Here it is made by other processes which need not be explained in this place. It is now an article of commerce and in large cities can be obtained in any quantity desired com- pressed in strong iron cylinders. OZONE. Occurrence. Along with the oxygen in the atmosphere several other so-called oxidizing substances are known to exist in small amount. One of these is a peculiar form of oxygen itself and is called ozone. It must be remembered, however, that the trace of this substance in the air is so minute that tests for its presence often fail. GENERAL CHEMISTRY. 41 History. The peculiar odor noticed in the neighbor- hood of a plate electrical machine in action was long ago remarked. In 1785 Van Marum called attention to the fact that the same odor is developed by the passage of electric sparks through pure oxygen, and showed that the gas so acted upon has the power of immediately tarnishing a clean surface of mercury. He found also that a decrease in the volume of oxygen taken follows. In 1801 Cruik- shank observed that in the electrolysis of water acidulated with sulphuric acid a peculiar odor is developed at the positive pole. Of the nature of the substance having this marked odor neither he nor Van Marum had any knowl- edge. It was reserved for the German chemist, Schoen- bein, to publish the first definite details touching the production and properties of this new body, which he did in 1839 and 1840. Schoenbein showed that by the pas- sage of electricity through air or oxygen, in the electrolysis of water acidulated with certain acids and in the oxidation of moist phosphorus the same substanceis formed, to which he gave the name ozone. The true composition of the gas was not recognized immediately. A number of important investigations, extending through fifteen years, were required to fully settle it to the satisfaction of all. Preparation. By special methods we can produce ozone in the laboratory, although not easily in large amount. It will suffice here to illustrate this fact by a very simple experiment. Ex. 47. Scrape the surface of a stick of phosphorus, about 5 Cm. long, under water, so as to expose the pure substance, free from coating or incrustation. Next pour some lukewarm water into a tall beaker or wide mouth bottle to a depth slightly less than the thickness of the scraped phosphorus. Then transfer the latter to the beaker or bottle and see that the surface is exposed to a slight extent above the water to the oxidizing action of the air. Cover the vessel with a piece of glass and allow it to stand some time; five minutes is usually sufficient. Meanwhile prepare a test for the ozone, and this can be done for the present instance in the following manner: Dissolve a small crystal of pure potassium iodide in distilled water and into this solution stir a little starch. Rub the starch with a glass rod to break up any lumps formed and then gradually heat to boiling This makes a paste in which the 42 GENERAL CHEMISTRY. potassium iodide exists dissolved. Into this paste dip some small strips of filter paper and then suspend these in the vessel in which the phos- phorus was left in contact with the moist air. If the experiment has been properly performed the paper will turn blue, indicating ozone formed. The following explanation must be given of the above experiment. Potassium iodide is a substance which is not decomposed either in the dry state or in solution by ordi- nary oxygen, but it is decomposed by other substances, among which is ozone, with liberation of one of its con- stituents iodine. This iodine forms a deep blue color with starch paste. If, therefore, we have a mixture of potassium iodide and starch paste and if from any cause this becomes blue, we know that some strong decomposing agent has acted on the compound, the potassium iodide, setting free its iodine. In the present case the paper turns blue because the air in the bottle contains a small amount of ozone formed by the action of the phosphorus on the moist air. The ordinary oxygen has not the power of producing this blue color. This is an illustration of one of the many decomposi- tions which the ozone gas can effect. It is especially active in breaking up organic matters, oxidizing or burn- ing them in a certain sense. Ozone in the atmosphere is supposed to have the action of a purifying agent in de- stroying decaying animal and vegetable matters. What is commonly called ozone test-paper is prepared by covering good book paper with a paste of Water 1,000 parts Starch 50 Potassium iodide 5 " The potassium iodide is dissolved in a small amount of the water in a porcelain dish. Into this solution the starch is stirred and rubbed with a pestle until it forms with the water a uniform creamy liquid. The remainder of the water is then added, and the whole is heated, with con- stant stirring, on a water-bath until the starch is converted into a smooth paste free from lumps. This hot paste is spread over white paper by means of a soft flat brush, and GENERAL CHEMISTRY. 43 then the paper is hung up to dry in an atmosphere free from oxidizing gases. The dried paper, cut into small pieces, may be kept almost indefinitely in stoppered bot- tles. When used as a test for ozone in the air, a small piece is moistened and hung up in the atmosphere in ques- tion. It is used as a test for other substances, as will be shown later. Larger quantities of ozone may be made in so-called ozone generators. These are forms of apparatus in which a current of oxygen may be passed between metallic plates connected with the terminals of an induction coil. The silent discharge across the intervening space converts a part of the gas into ozone. But the reaction is always far from complete, unless the product is absorbed by po- tassium iodide solution, or something else, as fast as formed. Ozone has been shown to be formed by the conden- sation of ordinary oxygen in a peculiar manner, which will be referred to later. In this condensation 3 volumes of oxygen yield 2 of ozone. At a temperature of 300 C. this condensed product is completely decomposed, common oxygen resulting. The oxidizing action of ozone is powerful, many organic substances being quickly destroyed by it. As a bleaching agent it is many times as strong as chlorine. In the older literature (since 1850) it was considered as the most powerful natural purifying agent in the atmosphere, but it is now generally admitted that most of the effects ascribed to ozone in the air are due to a related body, the peroxide of hydrogen, which will be described later, or to nitrous acid, which is present in small traces. HYDROGEN. This is a gaseous element of the highest importance from many standpoints. Occurrence. As a free substance it is found in nature in traces only, but is one of the common elements in com- 44 GENERAL CHEMISTRY. bination, constituting one-ninth by weight of water and an important fraction of most animal and vegetable sub- stances. History. When iron and zinc are dissolved in dilute acids an inflammable gas is evolved. This fact was observed by some of the alchemists, but received no expla- nation from them or their followers. Cavendish, in 1766, published an investigation of the subject in which he de- scribes the gas as inflammable air. Later he considered it as identical with pure phlogiston, because it was found capable of regenerating pure metals from the calces referred to under oxygen, and this view was held by many others. Following up his experiments, Cavendish found with con- siderable accuracy the amount of the gas which may be liberated from acids by given weights of several metals. In 1781 he found that water is composed of inflammable air and dephlogisticated air, but at the time he apparently failed to realize the importance of his discovery. In 1783 Lavoisier gave the first clear explanation of the composi- tion of water and proposed the name hydrogen for the inflammable air of Cavendish. About this time it was found that hydrogen may be obtained by the action of cer- tain metals on strong alkali solutions as well as on acids. Preparation. We are able to separate hydrogen from its compounds by many simple reactions. In illustration of these we will consider first the decomposition of water. It has been stated before that water is a compound of hydrogen and oxygen, and in separating them we must overcome the strong affinity which holds them together. This may be conveniently done bypassing a strong cur- rent of electricity through the water slightly acidulated with sulphuric acid. Certain metals brought in contact with water are also able to separate the hydrogen through the attraction they have for the oxygen. The following experiment will show how this separation may be effected at the ordinary temperature: Ex. 48. Drop a small piece of sodium, not larger than a pea, on the surface of water contained in an earthenware bowl. As the metal is GENERAL CHEMISTRY. 45 lighter than water it floats, but as it does so a decomposition goes on in- dicated by the escape of a gas with a sound reminding one of the escape of steam. On striking the water the sodium melts and assumes the globular form. It soon becomes evident that in the reaction between the two the sodium is worn away, as the globule grows small and finally disappears. While the gas is escaping bring a small flame in contact with it and observe that it ignites and burns readily with a yellow color. Very often the gas ignites spontaneously. In performing this experi- ment it sometimes happens that the sodium globule flies into small bits which are scattered in all directions. The face should not be held over the bowl, therefore, when making the test. Sodium is a light, silver white metal which is kept under rock oil. When a piece is taken out for this experi- ment it is cut to the proper size and wiped free from the oil by means of filter paper. The liberated gas may be collected in a bottle or test-tube and examined. To do this, fill the bottle or tube with water and invert it in the bowl in the usual manner. Then wrap a bit of clean dry sodium in filter paper or wire gauze, and by means of tongs bring the pellet so enclosed under the mouth of the bottle or tube, held for the purpose just below the surface of the water in the bowl. The decomposition takes place as before, but the gas ascends into the receptacle, and may be tested as explained later. The beginner is not advised to make this experiment. The yellow color of the flame is not characteristic of hydrogen, but of the vapor of sodium burned with it. Ex. 49. Repeat the last experiment, using a small piece of potas- sium instead of sodium. Observe the same precautions. The flame is now violet. Ill both of these experiments hydrogen gas is set free and a certain amount of water has been converted into something else. The character of this is disclosed by two simple tests. First, by means of a glass rod touch a little of the water to the tongue; a sharp caustic taste will be noticed. Then dip a piece of red litmus paper in it and notice the change of color to blue. This is evidence of the presence of alkali. The same evidence is given by the deep red color produced when a few drops of alcoholic phenol-phthalein solution are added to the water. It appears, therefore, that in this experiment with the water 46 GENERAL CHEMISTRY. hydrogen gas is liberated, while an alkali substance is formed. Sodium and potassium are not the only metals which decompose water in the cold, while at a high temperature the reaction is possible with many others, as is readily shown. The usual method of procuring hydrogen for ex- periment depends on the decomposition of some acid by means of a metal. In Ex. 29 it was shown that when dilute sulphuric acid is poured over zinc in a beaker the FIG. G. metal gradually dissolves with evolution of gas. This gas is hydrogen and the experiment may now be arranged to collect and test it. Ex. 50. Arrange a gas generating bottle as shown in the cut above. The bottle should hold about 250 Cc. and be closed with a doubly per- forated stopper. Through one perforation the stem of a funnel tube passes while the gas is led out through a "delivery tube" from the other. Put some granulated zinc in the bottle, insert the stopper with the funnel tube and delivery tube, add some water and then some dilute sulphuric acid so as to about one-third fill the bottle and cover the lower end of the funnel tube. An evolution of gas soon begins, and if the end of the delivery tube is brought under water the gas passes through and may be collected as in the case of oxygen. Therefore arrange several bottles for the collection of gas as there described, and after it has bubbled through the water a few minutes, allow them to GENERAL CHEMISTRY. 4? fill with the gas. When filled remove them by aid of a glass plate but keep the mouths of the bottles now down, because the gas is much lighter than air. When several bottles are filled their contents may be tested. Ex. 51. Lift one of the bottles from the table, mouth still down, and thrust up into it a lighted taper or splinter of burning wood. As the light goes into the bottle it is extinguished but a flame appears at the mouth of the bottle, from the ignited hydrogen. The gas, there- fore, burns, but the combustion on the end of the taper or splinter is checked, because, as shown before, oxygen is necessary for this and the bottle contains hydrogen. The hydrogen itself, at the mouth of the bot- tle, enters into combination with the oxygen of the air. Ex. 52. Invert one of the filled bottles, holding the mouth now up, remove the glass plate and bring a flame to the gas. It will ignite with a slight report and burn in the bottle, as the heavier oxygen of the air tends now to settle and mix with the hydrogen. In nearly all cases the flame shows some color. This is not characteristic of the hydrogen but of various impurities with it or on the surface of the glass. The flame from pure hydrogen has a slight blue tinge only. Ex. 53. The relative lightness of hydrogen may be readily shown by pouring it upward into a bottle filled with air. Use one of the bot- tles of hydrogen still standing on a glass plate. Lift it with one hand, and turn it so that its mouth is brought under and near that of the air- filled bottle of the same size. After a minute or two hold the upper bottle to a flame, when a sharp report shows the presence of hydrogen. The lower bottle still contains some of the gas, which can be shown in the same manner. It will be seen later that certain gases heavier than air can be poured downward from one vessel into another, as in the case of water. This is true of carbonic acid gas and chlorine, for instances. In the above experiment the light hydrogen ascends, and a certain amount of the air in the upper bottle is forced out by it. The displacement is not perfect, however, as was shown by the manner of the ex- plosion of the gaseous mixtures in the two bottles above. Hydrogen from Other Sources. While hydrogen is commonly prepared as just shown, it may be made by the action of certain metals on alkali solutions. Aluminum wire, for instance, decomposes a solution of potassium or sodium hydroxide very readily, especially when aided by heat. The reaction may be carried out in a test-tube, and 48 GENERAL CHEMISTRY. the character of the escaping gas determined. The same alkali solutions are decomposed by other metals also. These reactions have practical value, as in certain investigations it is desirable to liberate hydrogen without the use of an acid, and the alkali methods may then often be applied. Hydrogen is also easily liberated from water by the passage of the electric current, as intimated. This will be illustrated later by an experiment to determine the com- position of water. It has been shown that sodium and potassium decom- pose cold water with liberation of hydrogen. At a higher temperature the same decomposition may be effected by other metals. When steam is passed through an iron or FIG. 7. porcelain tube containing iron turnings, and heated to a very high temperature in a gas furnace, it is decomposed, the hydrogen being set free, while the oxygen remains in combination with the iron, forming an oxide of iron. This experiment is easily carried out in the apparatus illustrated by the above figure. Water is boiled in a flask to the left. The steam generated passes into the tube resting over a number of burners in the furnace, while the liberated hydrogen is collected in a jar beyond. By filling the tube with charcoal instead of with iron turnings, a somewhat analogous decomposition takes place. We obtain now hydrogen mixed with oxides of carbon, as the carbon com- bines with the oxygen of the water to form these bodies, which are gases. GENERAL CHEMISTRY, 49 The last decomposition is a very important one, as it is the basis of the process commonly followed in the man- ufacture of water gas, generally used at the present time. It will be fully described later. Diffusion of Hydrogen. Because of its extreme light- ness this gas is very suitable for showing an interesting property of all gases, viz. : that of diffusion. Two gases separated from each other by a porous partition a thin plate of plaster of Paris, for instance will in time mix with each other, as both pass through the porous sub- stance. The rates of diffusion or passage of the gases bear a close relation to their specific gravities or densities. It has been found that the velocity of diffusion is inversely proportional to the square root of the density of a gas. From this it would follow that hydrogen must diffuse 4 times as fast as oxygen and 3.8 times as fast as air, as the densities of the gases stand to each other in the relation, 1 : 16 : 14.45. A simple proof that hydrogen moves much more rapidly than air is given in the following experiment: Ex. 54. From a piece of glass tubing having an internal diameter of a centimeter or more, cut off a length of about twenty centimeters. Dip one end of this into some soft plaster of Paris, so as to take up a plug about one centimeter in thickness. Set the tube aside for this to harden, which will require some hours Then fill it with hydrogen gas by displacement of the air, and immediately stand it in upright position in a beaker of water in such a position that the open end is covered to a depth of several centimeters. It will be observed that the water ascends in the tube and finally reaches a position much above the level of the water in the beaker. It then recedes slowly and in time the level cor- responds to that outside. In the first stage of this experiment the water ascends because the hydrogen passes out through the porous plug much faster than the air can enter. The maximum posi- tion of the water is reached when the rates of motion in opposite directions are equal, after which the column of water falls because the volume of gas entering the tube is now greater than that leaving it. Many porous stones may be used to exhibit this phenomenon. 50 GENERAL CHEMISTRY. Reducing Power of Hydrogen. By this we under- stand the property which hydrogen possesses of abstract- ing oxygen from certain compounds, forming with it water. The term is used also in a broader sense but in this place the limited usage only will be considered. It will be shown later by experiment than when hydro- gen gas is passed over the oxides of copper or iron, heated to a high temperature, the oxygen is taken and the metal left in the free state. Other metallic oxides may be reduced in a similar manner. This reaction is one of the highest importance and is frequently employed in the laboratory for several purposes. An illustration will be given in the next section. An animal placed in an atmosphere of hydrogen would soon die, but this would follow from asphyxiation rather than from any poisonous property of the gas. In an atmos- phere of 4 parts of hydrogen and 1 part of oxygen animals live apparently as well as in ordinary air. Hydrogen is but slightly soluble in water, 1 volume of the latter dissolving of the gas 0.0193 volume at a tem- perature of 0C., and under a pressure of 760 Mm. of mercury. Several metals possess the power of absorbing hydrogen in considerable quantity. In the case of the metal palladium this power is very marked. The cold metal absorbs about 375 times its volume of hydrogen, while at a red heat nearly 1,000 volumes are absorbed. Under standard conditions 1 liter of pure hydrogen weighs 0.0900 Gm. It has been condensed to a colorless liquid at a temperature below 200C. with a pressure of 40 atmospheres. WATER. Occurrence. The student is familiar with the natural occurrence of water in the seas, lakes, rivers, etc. The purification of water by distillation has been referred to already, and further details of practical processes will be given later. History. In the preceding section it was explained GENERAL CHEMISTRY 51 that the exact composition of water was first suggested by the experiments of Cavendish, while Lavoisier's work proved the fact conclusively. Careful investigations un- dertaken by Gay Lussac and Humboldt were published in 1805, and these confirmed the work of Lavoisier and showed that exactly two volumes of hydrogen combine with 1 volume of oxygen to form water. Some experiments will now be given to illustrate methods of finding this ratio. Composition of Water. The presence of hydrogen in water was suggested by the experiment in which metal- lic sodium was used to decompose water. Other metals, as intimated, behave in a similar manner. At a high tem- perature, water is readily decomposed by iron turnings with liberation of hydrogen, which has been referred to already. A still more convenient method of decomposition is by means of the electric current, which will be experi- mentally shown. Water in absolutely pure condition is not a conductor of electricity and therefore in this form is not decomposed by it. But by the addition of a little acid to water it be- comes a moderately good conductor, and the current which may now be made to pass through effects decomposition. Acids are not the only substances which when added to water render it a conductor, but for our purpose they are the most convenient. If two plates or strips of thin plati- num foil, attached to the opposite poles of a galvanic bat- tery of several Bunsen or Leclanche" cells, be dipped in a beaker of acidulated water, gas bubbles will be seen to ascend from the surface of each plate. If these plates be supported beneath two tubes filled with water the gas bub- bles will pass up into them and displace the water. After a time the contents of each tube may be tested. If the tubes are so placed that they collect all the gas given off from each plate in a given time it will be noticed that one volume is almost exactly twice the other, and that the tests of the larger volume show it to be hydrogen, while the smaller volume gives the tests characteristic of oxygen. Several special forms of apparatus have been devised for 52 GENERAL CHEMISTRY this experiment of which the one now to be described is very convenient. Ex. 55. Arrange the apparatus as shown in the next figure (Fig. 8). As seen it consists essentially of a long U tube, the two limbs of which are closed on top by ground glass stopcocks. From the bottom of the bend a tube passes backward and then upward, ending finally in a wide bulb or reservoir. Inside of each limb of the U tube, just above the bend, there is a thin, platinum plate, which is connected with a short platinum wire passing through the glass and ending in a loop just outside. FJG. 8. By opening the stopcocks on the U tube, the apparatus may be filled with acidulated water, to this level, by pouring into the bulb tube. The stopcocks are then closed. (Water containing about 5 per cent of sulphuric acid is suitable for the purpose.) The apparatus is now ready for the actual experiment which is begun by attaching the copper wires from a good battery to the platinum loops described above. Almost immediately gas bubbles form on each plate and escape up into the tubes above. As the plates are situated above the bend of the U, the gases cannot mix on liberation, but must remain separate. In a few minutes it becomes apparent that the volume of the gas liberated from the plate connected with the negative pole of the battery, the pole connected with the zinc, is practically twice the volume liberated at the other pole. GENERAL CHEMISTRY. 53 Allow the experiment to continue until the larger volume fills the limb down to the platinum plate. It will be observed that as the gases collect the liquid is forced down and then up into the bulb tube, and further that the diminution of the volume of liquid itself is not great. In fact, not more than a small drop of the liquid undergoes decomposi- tion to form the relatively large gas volume. The nature of the gases in the two limbs may now be tested. To this end wipe off the tips above the stopcocks, and free them as far as pos- sible from liquid by means of bibulous paper, then light a small taper or splinter, hold it over the limb with the larger volume and carefully open the stopcock. The pressure of the liquid in the bulb tube will force the gas out, and this burns in a manner characteristic of hydrogen, on com- ing in contact with the small flame. Over the other limb hold a glow- ing splinter. On opening the stopcock the gas which streams out gives the behavior characteristic of oxygen. According to this experiment it would appear that water may be decomposed into two volumes of hydro- gen and one of oxygen. By repeating the above experiment carefully with ac- curate apparatus it can be shown that one volume is a lit- tle more than twice the other. The oxygen volume is rel- atively small because this gas is more soluble in the liquid than is the hydrogen, and also because a little of it is changed into ozone in the reaction. Both of these facts have been referred to already. If precautions are taken to avoid the production of ozone and if the solubility of the oxygen is diminished by working at a high temperature then the two volumes will be found to be liberated in the exact proportions of 2:1. As carried out, the last experimentcould not be regarded as conclusive, as not all of the liquid was decomposed. But the experiment may be repeated as often as desired with the water remaining after each test until its volume becomes quite small. The result of the electrolysis is the same in all cases. Further information concerning the composition of water is given by the following experiments. If, as suggested by the foregoing, water is composed of hydrogen and oxygen gases we should be able to produce water by the combination of these two substances. The experiment may be readily made and according to several plans. Ex. 56. First, we may make a very simple experiment with the 54 GENERAL CHEMISTRY. apparatus illustrated below. To the common hydrogen generator is attached a so-called drying tube, which contains some substance to absorb moisture from the gas. Calcium chloride is often used for the purpose. The generated gas after passing through the drying tube reaches the air by means of a narrow tube with a fine opening. The hydrogen gas, after having passed long enough to expel the air, is lighted at this opening and burns with a small hot flame. If now a cold dry beaker be held, mouth down, over this flame a deposit of moisture immediately collects on the cold surface of the glass. This must be formed by a union of the hydrogen with something in the air. FIG. 9. That this substance from the air is the oxygen may be shown by trial with the pure gas. Given volumes of the two gases may be combined directly by aid of an electric spark and the result noted, when it will be found that exactly two volumes of hydrogen unite with one volume of oxygen. For our purpose it will prove more convenient to make the determination in another manner, somewhat less direct, but equally con- clusive. Instead of combining the hydrogen with pure oxygen we can allow hydrogen gas to act on some sub- GENERAL CHEMISTRY. 55 stance containing oxygen in combination, under such con- ditions that the substance will decompose, giving up its oxygen to the hydrogen. In such a case hydrogen acts as a reducing agent, as was explained some pages back, and the oxides of copper and iron are good illustrations of sub- stances which may be decomposed in this manner. The oxide of copper is a compound easily made in condition of great purity so that there is no doubt regarding its exact composition and the proportion of oxygen it contains. The experiment which the student is able to make with this substance is a qualitative one, but the modifications neces- sary to make it quantitative will be suggested. FIG. 10. Ex. 57. Arrange the apparatus as shown in the above figure. It consists of a hydrogen generator and a tube following it filled with cal- cium chloride to dry the gas. Then follows a piece of hard glass tub- ing, plain or furnished with a bulb, to contain the pure dry copper oxide to be decomposed. After the tube of hard glass comes a bent, or U tube, to catch the products of the reaction to be described. This U tube is immersed in a beaker of cold water, and in turn is joined to a second small calcium chloride tube. The connections must all be made with perfect, sound corks, which fit with absolute accuracy. The appa- ratus being ready, and the hard glass tube charged with 10 to 20 Gm. of the copper oxide, we begin the experiment by generating hydrogen in the usual manner by the action of dilute sulphuric acid on zinc. The 56 GENERAL CHEMISTRY. gas, given off in the moist condition, is dried as it passes through the calcium chloride tube and then enters the combustion tube, from which it displaces the air. After the gas has passed some minutes through the apparatus the flame of a Bunsen burner is applied to the combustion tube immediately under the portion of the oxide of copper nearest the generator. The heat must be very carefully applied at first to avoid cracking the hard tube, which can be best prevented by having the flame low and by moving it to and fro along the tube. After a few minutes stronger heat may be applied . It will soon be noticed that vapor ascends from the black mass and that it is pushed forward by the pressure of the hydrogen toward the U tube, and also that in a short time the heated oxide above the flame glows as if on fire. We have here the stage of active reduction. It becomes presently evident that in this very hot part of the tube the oxide of copper has been converted into bright metallic copper, the color of which is very distinct. When the reaction here is complete the flame is moved onward toward the U tube and the operation continued until the whole of the black oxide has undergone decomposition. The vapors given off suggest the formation of water, but this can be further tested in the U tube. These vapors tend to con- dense in the cooler part of the combustion tube itself, but the condensed product can be dislodged by careful application of the lamp heat and completely driven over into the U tube in the beaker of water. A small portion of the vapor passes through the U tube, but is intercepted by the calcium chloride tube beyond. After the completion of the reduction, that is, when the black oxide of copper has disappeared, leaving only bright metal, the heat is withdrawn, leaving the tube to cool down with the current of hydrogen still passing. Then it is taken apart and the contents of the U tube examined. The proper tests, which the student is not prepared to make at this stage of his work, however, show that the condensed liquid is water, and nothing else. It has evidently been formed by the union of the hydrogen gas with the oxygen taken, as needed, from the copper oxide. It sometimes happens in this experiment that the water collected is slightly colored. This is due to the action of heat on the corks at each end of the combustion tube, and can be avoided by careful manipulation. The oxide of copper should be pure and free from any kind of organic dust, and should occupy the central portion of the combus- tion tube. This must be quite clean and dry at the begin- ning of the experiment. The oxide of copper may be held in proper place by means of a loose tuft of clean glass wool at each end. To make the experiment perfect in all details would necessitate the complication of the apparatus to a degree which would render it unfit for our present purpose. As GENERAL CHEMISTRY. 57 constructed and manipulated above it is sufficient to show conclusively that only hydrogen and oxygen are concerned in the formation of water. By the introduction of certain modifications the experiment may be made a quantitative one. It is simply necessary to provide more perfect ap- paratus to absorb the water formed, and weigh this accu- rately before the experiment. It is weighed again at the end, the increase in weight showing exactly the amount of water produced. The oxide of copper must be weighed before the experiment, and the residue left after it. The loss of weight here corresponds to the oxygen given to the hydrogen. The experiment, therefore, gives us the weight of oxygen in a determined weight of water; the difference between these two weights represents the hydrogen, be- cause nothing else has been used in the work. This is a fundamental experiment upon which many skilled chemists have spent a great deal of time, for the purpose of determining with the highest possible accuracy the exact ratio in which hydrogen and oxygen combine. This knowledge, as will appear later, is of great scientific importance, as much depends on it. A very satisfactory method of showing the direct union of the two gases in the proportion mentioned, is by means of an explosion of a mixture of the gases in an apparatus known as an eudiometer, shown in the next figure. This apparatus consists essentially of a long, accurately grad- uated glass tube, the divisions being usually in millimeters. Near the closed end of the tube two pieces of platinum wire are passed through the glass and sealed in so that they nearly, but not quite, touch inside the tube. The eudiometer is filled with mercury, inverted in a reservoir of mercury, b, and clamped in position. A certain volume of pure hydrogen is now allowed to enter the tube and is accurately measured, the tube being brought to a vertical position for this purpose, and for all subsequent measurements. The illustration shows the tube in an inclined position, into which it is brought for convenience in filling. A volume of pure oxygen over half as great is next added and the new mixed volume accurately noted. The 58 GENERAL CHEMISTRY. open end of the eudiometer is pressed down on a piece of rubber under the mercury and then the tube is firmly clamped. A spark from an induction coil is caused to jump between the wires within the tube. This produces an explosion in the gaseous mixture, and when the pres- sure on the rubber plate is released it will be noticed that mercury rushes up into the tube. After a time the remain- ing gas volume is accurately measured. A subsequent test shows that it is oxygen. If this volume is subtracted from the oxygen volume taken it will be seen that the FIG. 11. oxygen actually used is just one-half of the hydrogen volume taken. The droplet of water formed in the ex- plosion is so small that its volume may be neglected in comparison with the gas volumes concerned. In measuring the gases in this experiment certain pre- cautions must be taken which are fully explained in a fol- lowing chapter. In all such cases temperature and pressure must be accurately observed so that by reduction all the volumes may be compared under the same condi- tions. GENERAL CHEMISTRY. 59 Physical Properties of Water. Pure water has a constant freezing point and constant boiling point under constant conditions of pressure. It is a poor conductor of heat and is practically a nonconductor of electricity. When water at the ordinary temperature is heated it expands pretty regularly until its boiling point is reached, and by further addition of heat is converted into steam, the volume of which is about 1,700 times that of the water. When cooled a contraction of volume follows until the temperature of 4 C. is reached. At this temperature, or more accurately, 3.95, water reaches its smallest volume for a given weight, and therefore, its maximum density. When further cooled it expands slightly until the freezing point is reached. On conversion into ice a marked expan- sion takes place, 100 volumes of water at yielding 109.1 volumes of ice. Capacity for Heat. Unit of Heat. It is a matter of common experienc~that water "absorbs" a large amount of heat, practical application of which power is made in the hot water system of heating houses. In order to measure any amount of heat under consideration phys- icists have adopted what is known as a unit of heat. This may be arbitrarily defined as the amount of heat necessary to raise the temperature of a gram of pure water through one centigrade degree. To be scientifically correct this defini- tion requires a slight modification or qualification, but it is sufficiently close for the present purpose. Seme interest- ing facts have been brought out in studying the absorption of heat by water in its three forms. To illustrate these let us start with a gram of ice at centigrade, that is, at its melting point. To convert this into water at the same temperature, a relatively large amount of heat must be applied. It has been found that 79.5 units must be added to accomplish this. In other words, as much heat must be absorbed by the ice in melting as would be sufficient to warm 79.5 grams of water one degree, or one gram of water 79.5 degrees. This absorbed heat is usually spoken of as latent heat, because it is lost or hidden as far as any 60 GENERAL CHEMISTRY. thermometric observation is concerned. It is used up, however, in doing a certain kind of work on the ice, that is in changing its physical condition. If we continue the addition of heat after the ice has been melted the effect now becomes visible in the increase of the temperature of the water. For each unit of heat added the temperature of the water rises one degree centigrade. Finally, on addi- tion of 100 units of heat the water begins to boil, and a centigrade thermometer immersed in it marks 100 degrees. Supposing the water in a vessel under constant pressure, that of the atmosphere, it is now observed that further addition of heat produces no corresponding elevation of temperature. It is a well-known fact that a thermometer immersed in pure boiling water registers a constant tem- perature. The heat applied is, therefore, again rendered latent as in the case of the conversion of ice into water. Now it is used up in doing a new kind of work, the conversion of the water into steam at the same tempera- ture. Accurate experiments have shown that about 536 units of heat are required to convert a gram of water at a temperature of 100 into steam at the same temperature. When a gram of steam at this temperature condenses it gives out the 536 units. Many of thetechnical applications of water and steam depend on these remarkable properties. Many other liquid and solid substances may be brought to a higher temperature than is possible with water, but for a given range of temperature not one of them absorbs nearly as much heat. A hot water bag filled with water at 60 C. will give out, as it cools down, nearly ten times as much heat as would a mass of iron of the same weight at the same temperature. The specific heat of iron and all other solid and liquid substances is much below that of water. This matter will be taken up again. Boiling Point of Water and Vapor Tension. By definition water is said to boil at a temperature of 100 centigrade. But this temperature of ebullition depends on the pressure on the surface of the water, usually that of the air. If the atmospheric pressure is normal, that is, if it is equivalent to the pressure of a column of mercury 760 GENERAL CHEMISTRY. 61 Mm. in height the boiling point is constant at 100. At pressures below this the boiling point is below 100, and at higher pressures it is above. In any case water boils when the tension of the vapor which it gives off is equal to the pressure of the atmosphere. When water is heated in a confined space, as in a boiler furnished with a weighted valve, the temperature of ebulli- tion may become very high. Corresponding to any given pressure on the surface of the water there is a definite temperature of ebullition. Even at very low temperatures the vapor given off from water possesses a certain tension. This is shown in the table below where the tensions cor- responding to certain high as well as low pressures are given. The tensions are expressed in millimeters of mer- cury. Temp. Deg. Tension. Temp. Deg. Tension. Temp Deg. Tension. Temp. Deg. Tension. 1 4.909 18 15.330 35 41.78 120 1491 2 5.272 19 16.319 40 54.87 125 1744 3 5.658 20 17.363 45 71.36 130 2030 4 6.069 21 18.466 50 91.98 135 2354 5 6.507 22 19.630 55 117.52 140 2718 6 C.672 23 20.858 60 148.88 145 3126 7 7.466 24 22.152 65 187.10 150 3581 8 7.991 25 23.517 70 233.31 155 4089 9 8.548 26 24.956 75 288.76 160 4652 10 9.140 27 26.471 80 354.87 165 5275 11 9.767 28 28.065 85 433.19 170 5962 12 10.432 29 29.744 90 525.47 175 6717 13 11.137 30 31.510 95 633.66 180 7546 14 11.884 31 33.37 100 760.00 185 8453 15 12.674 32 35.32 105 906. 190 9443 16 13.510 33 37.37 110 1075. 195 10510 17 14.395 34 39.52 115 1269. 200 11689 From the table it is seen that at a temperature of 150 the tension of the vapor on the containing vessel is 3,581 millimeters of mercury, or over 4.7 atmospheres. Water as a Standard Substance. Because of the fact that water is everywhere abundant and readily ob- 62 GENERAL CHEMISTRY. tained in pure condition, it is well suited for use as a standard substance. It is employed in the definition of a unit of heat as mentioned above. In the construction of thermometer scales it is employed in determining two fixed points. In the centigrade or Celsius scale one fixed point is called the zero point, and represents the tempera- ture at which water freezes. The other fixed point is marked 100 and represents the boiling point of water at the normal pressure of the air at sea level. By definition a gram is taken as the weight of a cubic centimeter of pure water at a temperature of 4 C. at the latitude of Paris. By specific gravity or specific weight we understand the ratio of the weight of a given volume of substance to the weight of the same volume of pure water as the standard. In many other cases water is taken as one of the elements in comparison, but the above illustrations are sufficient. Solvent Action of Water. Water is the best single solvent known, although for practical purposes many sub- stances may be considered as insoluble in it. Of those which do dissolve some are much more soluble than others. One volume of water dissolves 0.041 volume of oxygen at a temperature of C. and a pressure of 760 Mm., but under the same conditions it dissolves nearly 80 volumes of sulphurous oxide and about 500 vol- umes of hydrochloric acid gas. Some mineral substances, as gypsum, are but slightly soluble while others, salt and saltpeter for illustrations, dissolve very largely. With in- crease of temperature there is in most cases a marked increase in the degree of solubility of bodies. Even such hard substances as the glass of our beakers and test-tubes dissolve to a very slight extent by long contact with boiling water. This fact is of great importance in some branches of chemical analysis. Natural Waters. All natural water comes to the earth in the form of rain, and in this condition it is nearly pure, containing not much more than traces of dissolved gases. This is especially true of the water collected at the end of a shower, that which falls first serving to GENERAL CHEMISTRY. 63 down the dust suspended in the air. On reaching the earth the character of the water is very speedily modified by the mineral substances which, by virtue of its solvent action, it takes up. As the rain descends it dissolves a little carbon dioxide and thus becomes a very weak solu- tion of carbonic acid. This acid aids in the solution of limestone and other substances from the soil, and in this way the water becomes hard. Hard waters are those which hold in solution relatively large amounts of certain mineral substances, principally salts of calcium and mag- nesium. If the soil on which the rain falls, and through which it filters or percolates, is free practically from these soluble mineral constituents, the water appearing later as a spring or brook is called soft water. Some natural waters contain not over 50 milligrams of dissolved substances in one liter, while others contain in a liter many grams. Purification of Waters. Natural waters contain substances dissolved and substances in suspension. Both may be objectionable for certain purposes, and before use water must often be freed from them. The highest degree of purification is ordinarily accomplished by distillation, the condensed steam being free from the dissolved solids and suspended matters originally present. In many cases purification extends only to a mechanical separation of suspended matters, which is accomplished by filtration through sand, charcoal or other porous substance. Fi- nally, water is frequently treated on the large scale with something that produces in it a bulky precipitate which in settling carries down practically everything in suspension. After such precipitation, by milk of lime or alum, the water is allowed to settle thoroughly, or clarification is hastened by filtration. HYDROGEN DIOXIDE. Hydrogen and oxygen combine in two proportions. In one case water is the result and in the other the body known as hydrogen dioxide or peroxide. 64 GENERAL C HEM IS TR Y. Occurrence. This substance occurs in some atmos- pheres in small traces, being produced by several natural agencies. In the air it is frequently confounded with ozone, as the two bodies are much alike in their behavior with reagents. Preparation. Hydrogen dioxide can easily be made in the laboratory as illustrated by the next experiment. Its preparation depends on the decomposition of a solid sub- stance, usually barium dioxide, by means of sulphuric acid. Ex. 58. Take about 10 grams of barium dioxide in a beaker, moisten it with water and allow the mixture to stand half an hour or longer, with occasional stirring. Then add about 20 cubic centimeters of dilute sulphuric acid (1 vol. of acid to 10 of water), stir well and af- ter a few minutes filter. The acid acts on the barium dioxide in the dry state very slowly, but if the latter has been previously hydrated by the action of water the decomposition is much more rapid. Insoluble barium sulphate and soluble liquid hydrogen dioxide result. The former is separated by filtration while the solution of the latter serves for tests. Properties. As usually made, hydrogen dioxide is largely mixed with water, but in the pure state it is a'thick- ish liquid with a specific gravity of 1.45. In this condi- tion it is not stable, but decomposes at the ordinary temperature into water and oxygen gas. For this reason the substance is always kept in very dilute condition, and experience has shown that its stability is increased by having a little free sulphuric acid in the solution. Even in dilute solution it is characterized by its strong oxidizing properties, and the numerous uses it has found in the arts and medicine depend on this fact. Tests. Some of the reactions of hydrogen dioxide may be shown by these experiments: Ex. 59. To a solution of the dioxide add some solution of potas- sium iodide. A decomposition of this compound takes place, with liberation of iodine, as is shown by the brown color of the liquid. If a little starch paste is added, it turns blue. The solution of the dioxide used in this test should be nearly neutral, that is, free from anything greater than traces of acid. This test, it will be observed, is very simi- lar to that for ozone, given earlier. GENERAL CHEMISTRY. 65 Ex. 60. Another very interesting decomposition is shown as fol- lows: Acidify the solution of the dioxide with dilute sulphuric acid and add to it a dilute aqueous solution of potassium permanganate, a few drops at a time. This latter solution has a deep purple color and as the drops fall into the dioxide liquid the color disappears, while bubbles of oxygen gas escape. On addition of a sufficient amount of the perman- ganate solution the purple color persists. The chemical reaction taking place here is somewhat complicated and cannot be explained in this stage of the work. The two experiments just given illustrate the marked property of producing decompositions possessed by the dioxide. This property depends on the fact, to be more fully explained later, that part of the oxygen united to hydrogen in the compound is only loosely held. It is very easily liberated and performs highly characteristic oxida- tion changes in consequence. Hydrogen dioxide is used to some extent in laboratories as a reagent but its most common applications are as a bleaching agent and in medicine. CHAPTER III. CHLORINE AND HYDROCHLORIC ACID THEO- RETICAL CONSIDERATIONS. WE COME now to the consideration of some very important substances which never occur in the free state in nature, but in many compounds are widely distributed. CHLORINE. Occurrence. This is an exceedingly abundant sub- stance in combination, being one of the constituents of common salt. The other constituent is a body called sodium. It is found also in many compounds somewhat similar to salt, all of which are called chlorides. Salt is known as sodium chloride, while the next most abundant chloride is potassium chloride. History. Chlorine was first prepared by the great Swedish chemist, Scheele, in 1774, and by a method which is still commonly employed for the purpose, viz., by the action of black oxide of manganese on hydrochloric acid. Scheele did not, however, recognize the true nature of the substance, and it remained for Humphrey Davy, in 1810, to supply this information and propose the name now given to the body. Preparation. All methods of preparing chlorine de- pend on the decomposition of some of the chlorides. Usu- ally we employ sodium chloride or hydrogen chloride, called also hydrochloric acid. The preparation by the use of hydrochloric acid will be illustrated first. This sub- GENERAL CHEMISTRY. 67 stance, as will appear later, is a compound of hydrogen and chlorine, and our problem is to separate one from the other. The next experiment will show how this can be done. Ex. 61. In a flask holding 300 Cc. or more, take about 50 Gm. of manganese dioxide, the substance already used with potassium chlorate in the preparation of oxygen. Pour over the dioxide about 200 Cc. of commercial strong hydrochloric acid. Close the flask with a stopper FIG. 12. having two perforations. Through one of these a funnel tube passes, the lower end dipping into the acid. A delivery tube passes out through the other perforation. This extends up about 6 or 8 Cm., and there is bent at right angles, the horizontal portion following having about the same length. To this, by means of a rubber connection, is joined a second bent glass tube, the longer limb of which has a length of about 20 Cm. The flask is supported on a sand-bath, as shown in the above figure, while the delivery tube extends down into an empty dry bottle of 200 to 300 Cc. capacity. When the flask is charged and properly mounted apply a gentle heat to the sand-bath. This hastens the action of the acid on the manganese dioxide. A greenish yellow gas soon fills 68 GENERAL CHEMISTRY. the flask and passes over into the dry bottle. The delivery tube should enter the bottle through a piece of perforated cardboard, in order to prevent, as far as possible, the escape of the gas while the bottle is fill- ing. As chlorine gas is much heavier than the air, it can be collected in this manner by displacement, the air being driven up out of the bottle. When the bottle appears to be quite full of the gas remove it and put a second in its place. Collect four or five bottles of the gas. Each bottle as removed from the generator must be covered by a glass plate. After collecting the desired amount of the gas, replace the dry bottles by one containing water. Continue the application of heat and add more hydrochloric acid if necessary. Chlorine gas is somewhat soluble in water, and in this manner a solution is obtained which is used in a following experiment. All experiments with chlorine must be per- formed in a fume closet. Save the liquid in the generator. We have now several bottles of the gas, of which the greenish yellow color is characteristic. The odor also is marked and disagreeable in the extreme. The student must avoid, as far as possible, inhaling it, as it is exceed- ingly irritating to the air passages. In large amounts it is even very dangerous. For this reason the direction is given to carry out all experiments in a fume closet where there is sufficient draught to carry off any escaping gas. Having become aware of the more prominent features of the substance, the student may determine some of its prop- erties by experiment. Ex. 62. Chlorine is a strong bleaching agent, which maybe shown as follows: Moisten a strip of colored calico in water and suspend it in one of the bottles of the gas, replacing the glass plate after introducing the fabric. In time the color fades through the destructive action of the chlorine. In the dry state the gas is practically without action, hence the direction given to moisten the calico. In the same bottle pour a little solution of indigo. If the bleaching of the calico has not removed all the chlorine the indigo color will be destroyed likewise. This important property of the gas is utilized on the large scale in the wholesale bleaching of many articles. The action however is often destructive of the organic fiber itself. Ex. 63. Chlorine has a marked affinity for many metals as well as for organic colors, and this may be illustrated by a very simple test. Remove the glass plate from one of the bottles, filled as above described, and put in its place a piece of wire gauze. Through this sift into the gas a little powdered antimony. The fine metal particles ignite as they fall in a shower through the gas and burn brightly, forming a chloride GENERAL CHEMISTRY. 69 of antimony. Many other metals burn equally as well if they are taken in the form of fine wire or foil. Antimony in powder exposes a large surface to the action of the gas. Chlorine was prepared above by decomposing its compound with hydrogen. It may readily be made to com- bine with hydrogen, again reproducing the acid. If equal volumes of the two gases be mixed in the dark and then exposed to the light, combustion of ten follows with a sharp explosion. The student is not advised to attempt this experi- ment as it is dangerous, unless carried out with certain pre- cautions which need not be described here. The same affinity of hydrogen for chlorine may be shown in an- other manner without risk as will now be explained. Many substances are known which consist of hydrogen and carbon only. These bodies are called hydrocar- bons. It has been found by experiment that chlorine gas is able to decompose a number of these substances, combining with the hydrogen to produce hydrochloric acid, while the carbon is set free as a fine black powder or soot. Among the hydrocarbons which exhibit this behav- ior, turpentine oil serves our purpose best, and wifl be employed therefore in our test. Ex. 64. Pour some oil of turpentine in a test-tube, and in it im- merse a strip of filter paper which has been twisted in the form of a taper. After withdrawing the paper press against it a second piece of dry filter paper, in order to absorb the excess of the oil taken up. Now remove the glass plate from a chlorine bottle and into it dip the taper, which in a few seconds darkens and finally burns with a very sooty flame, sending up a large volume of smoke. This consists of the liberated car- bon, while hydrochloric acid vapors are formed, as can be shown by proper tests. For the success of this experiment the chlorine gas must be practically free from air, and the paper must not be wet but only moist from the action of the turpentine. , The same general behavior is shown by burning a small wax candle, or even a small jet of illuminating gas from a bent glass tube, in a bottle of the gas. Both the wax and illuminating gas contain hydrogen in combination with carbon. A combustion of either of these substances, begun in the air, continues in chlorine with a very sooty flame. A small portion of the solution of chlorine water pre- pared above may be used for a bleaching test, but the 70 GENERAL CHEMISTRY. larger portion should be preserved in a stoppered bottle, kept in a dark place, for several tests to follow. Ex. 65. Prove that the chlorine water bleaches as does the gas, by immersing a piece of calico in it, or by pouring in some solution of indigo or litmus. The method given above serves very well for making small quantities of chlorine for experimental purposes, but foi the preparation of larger quantities the following proc- ess is much better. . The gas is generated in a large flask arranged as before, but from a mixture of salt, manganes'e dioxide and sul- phuric acid. Mix about equal weights of the salt and the dioxide ; pour this into the flask, and then through the funnel tube add gradually to the mixture about twice its weight of sulphuric acid, previously diluted with an equal weight of water. The flask is slowly heated on a sand-bath, and the gas is given off gradually. If the gas is to be used for making pure chlorine water it should be led through a wash bottle first. In its sim- plest form this may consist of a wide mouth bottle holding about 250 Cc., closed with a stopper with two openings. Through one of these a tube from the generator passes and dips beneath the surface of water, which about half fills the bottle. A second tube passes up from the under surface of the stopper and bends over, to lead the gas to water or to collecting bottles, as desired. By this arrangement the gas is washed by bubbling through the water in the wash bottle. It loses small amounts of hydrochloric acid and other impurities carried over from the generator, as these are more soluble in water than chlorine is. At 20 C. 1 volume of water dissolves 2.15 volumes of chlorine, and the solution so made can be kept a long time in a glass stoppered bottle in the dark. But exposed to light, gradual decomposition takes place, the hydrogen of the water combining with the chlorine to form hydrochlo- ric acid, while the oxygen is liberated. In bright sunlight the action is comparatively rapid, as can very easily be shown by experiment. Ex. 66. Fill a liter bottle with distilled water and pass chlorine GENERAL CHEMISTRY. 71 gas into it to complete saturation. Close the bottle then with a stopper, so as to exclude all air, and invert it in a jar containing strong chlorine water. Remove the stopper and stand the jar, with inverted bottle, in a window exposed to sunlight. In a short time gas bubbles will be seen to ascend through the liquid in the bottle. With average sunlight several days may elapse before the reaction is complete. The gas col- lected over the water in the bottle may now be tested. Insert the stop- per in the bottle, bring it to the upright position, then withdraw the stopper and apply the usual test for oxygen by means of a glowing splinter. The reaction is distinct and characteristic. If the bottle is allowed to stand long enough in the light the green color and the odor of the chlorine disappear, while the sharp, hydrochloric acid odor can be noticed. Other Methods of Preparation. Chlorine can be liberated by several other reactions, some of which have technical importance. One of these may be illustrated here by a brief experiment. Ex. 67. Take a gram or two of powdered potassium dichromate in a test-tube and pour over it a few cubic centimeters of strong commer- cial hydrochloric acid. Apply heat. Decomposition of the hydro- chloric acid soon takes place as shown by the appearance of greenish yellow fumes in the test-tube; that these consist of chlorine may be inferred from the color, odor and bleaching action easily determined. An explanation of this reaction will be given later. A strong solution of hydrochloric acid may be decom- posed by electricity in apparatus similar to that employed for the electrolysis of water. Several applications are made of this fact. On the large scale chlorine is liberated in quantity and cheaply by what is known as the Deacon process, from the name of the discoverer. In this process a stream of hydro- chloric acid gas is blown along with a stream of air through a series of heated tubes containing crushed brick impreg- nated with copper sulphate. The oxygen of the air takes the hydrogen of the acid and chlorine is left free, mixed with the nitrogen of the air. The crude chlorine so made is suitable for the production of bleaching powder and other products. General Tests for Chlorine. The common proper- ties of chlorine are so marked that they afford easy means of recognition. Like ozone and hydrogen peroxide, chlo- 73 GENERAL CHEMISI'RY. rine is able to decompose potassium iodide, and hence the so-called ozone test-paper, described in an earlier section, serves also as a chlorine test, when applied in moist con- dition to the gas supposed to contain or consist of chlo- rine. The methods by which chlorine may be recognized when mixed with other gases are described in Qualitative Analysis. Physical Properties. As stated above, chlorine is somewhat soluble in water. One volume of water at 20 dissolves about 2.15 volumes of chlorine. With ice water it forms a crystalline compound. Under strong pressure gaseous chlorine may be condensed to a liquid having a specific gravity of 1.33. This liquid is now an article of commerce. One liter of chlorine gas, at and under a pressure of 760 Mm., weighs 3.18 grams. Uses of Chlorine. While chlorine has many applica- tions on the large scale, it has also some in the laboratory. The gas itself is frequently used and also the solution, or chlorine water. It was directed above to save some of this solution, and with it several experiments will be made, as explained a few pages in advance. HYDROCHLORIC ACID. History. This acid was known in crude form to the Arabian chemists and was made from salt and green vitriol in the 15th century. About the middle of the 17th century it was first made by a process like that employed to day. It has been already intimated that this substance may be formed by the direct union of chlorine with hydrogen, also by the action of chlorine on water in sunlight. It is usually prepared, however, by the decomposition of a chloride by means of sulphuric acid. The cheapest chloride known is sodium chloride, or common salt, and hence this substance is nearly always employed in the preparation. The reaction may be carried out very easily as a laboratory experiment, by a method now to be given. GENERAL CHEMISTRY. 7:* Ex. 68. Arrange apparatus as shown by the illustration. The flask, to the left, on a sand-bath has a capacity of about 500 Cc. It is charged with about 50 grams of common salt, and is closed by a stopper with two perforations, through one of which passes a funnel tube leading nearly to the bottom of the flask. Through the other perforation a delivery tube passes, and this ends finally in a Woulfe bottle, half filled with water, but the delivery tube must not dip beneath the surface of the water here. Another tube leads from this first Woulfe bottle into a second, likewise half filled with water. In this case the tube dips beneath the surface of the water. From the second Woulfe bottle a tube leads to a flask of water. Each one of the Woulfe bottles has three openings. Through one of these, in each case, a so-called safety tube passes and dips into the water. The object of these safety tubes is to provide for easy communication with the air in case the pressure of gas in the generating flask should suddenly diminish FIG. 13. and the lower end of the funnel tube should be closed by the materials around it. When the apparatus is in order, the flask containing the salt, as explained, pour in about 50 Cc. of strong sulphuric acid through the funnel tube, a little at a time. Immediately a very lively reaction begins. The mass in the flask froths and rises, while gas bubbles escape through the water in the second Woulfe bottle and the adjoining flask. This is air being expelled. The hydrochloric acid formed by the action of the sulphuric acid is now seen to enter the first Woulfe bottle and pass down from the end of the delivery tube to the water and mix with it. The remarkable affinity of the water and gas is illustrated by this. After a time the flow of gas from the generator lessens and then heat should be applied to the sand-bath. At the beginning this is not necessary as the two substances react on each other in the cold. After a time the action 74 GENERAL CHEMISTRY. in the flask ceases and no more gas passes over into the Woulfe bottles. These are detached from the generator and their contents tested. Ex. 69. Begin the tests by taking equal volumes of liquid, as a test-tube full, from the two Woulfe bottles and the small flask. Pour the contents of the test-tubes into three small clean beakers, and add to each a few drops of solution of silver nitrate. In the liquid from the first Woulfe bottle a heavy, curdy white precipitate forms, in that from the second the amount of precipitate is much less, while in the case of the liquid from the end flask an opalescence only may result. Ex. 70. Remove now three more equal portions as before, trans- fer them to clean beakers and add to each five drops of a weak alcoholic solution of phenol-phthalein, which may cause a faint opalescence of no consequence. Have at hand a dilute solution of sodium hydroxide (caustic soda) and add this gradually to the contents of each beaker, beginning with that from the small flask. One or two drops of sodium hydroxide solution may be sufficient to impart a red color to the liquid, and this indicates that the acid present has been fully neutralized by the solution added, which is an alkali. The phenol-phthalein is a sub- stance which is turned bright red by alkaline solutions and hence it is employed here to show a change from acid to alkaline condition. It will be found that to neutralize the acid from the second Woulfe bottle more of the soda solution must be used, while for the liquid from the first Woulfe bottle a very large volume, relatively, of the alkali solution is necessary. It appears from this test that in the experiment most of the hydrochloric acid generated remains in the first Woulfe bottle. The last two experiments show certain important prop- erties of hydrochloric acid. It gives a precipitate with solution of silver nitrate, it neutralizes a strong alkali solution, and it is very soluble in water. We know this last to be true because both the silver and the alkali tests show that the most of the acid is in the water of the first bottle. That the acid is a gaseous body is indicated by the manner in which it is liberated in the beginning of the process when the sulphuric acid is first added to the salt. Without application of heat, it was seen that something passed over from the generating flask into the collecting bottle and this was evidently a gas. It has been found by experiment that at C. 1 vol- ume of water absorbs very nearly 500 volumes of the gas. The common hydrochloric acid which we handle in liquid form is merely an aqueous solution of the gas, containing from 25 to 40 per cent by weight of the real acid. For many purposes the strong acid is diluted, before use, with GENERAL CHEMISTRY. 75 more water. With about 10 cubic centimeters of the strong acid solution in the first Woulfe bottle make the following experiment. Ex. 71. Transfer the acid to a small heaker and add a few drops of the phenol-phthalein indicator. Then pour in very gradually, as before, some caustic soda solution until a red color just appears. Now, by means of a glass rod, add a drop or two more of the acid, or sufficient to discharge the color. This yields a very nearly neutral solution, the acid being only slightly in excess. Pour the liquid into a clean porcelain evaporating dish and boil it down to dryness. If during the evaporation the red color returns it shows that insufficient acid has been added and a drop or two more may be mixed with the liquid. When the evapora- tion is complete heat strongly a few minutes longer, allow to cool and observe the taste. We have here common salt, similar to that decom- posed in the large flask in the reaction with sulphuric acid. In this series of experiments we have illustrations of some very important chemical reactions. In the making of hydrochloric acid we decompose salt, which is a com- pound body, and secure two new substances. One of these is the gas hydrochloric acid which distills over, while the other is a solid substance, left in the generating flask, and is known as sodium sulphate. When it is seen that in neutralizing the hydrochloric acid with the soda solution we reproduce salt it becomes evident that the soda adds in the last experiment that which the sulphuric acid must have separated in the first. Hydrochloric acid on the large scale is often produced as a by-product in the manufacture of alkali from sodium or potassium chloride, as will be explained later. Sodium sulphate and potassium sul- phate, formed as above illustrated, are converted into sodium and potassium carbonates. Experiment shows that this acid is very active in the solution or decomposition of many bodies. It dissolves iron, zinc and several other metals forming chlorides, its hydrogen being liberated. It dissolves marble and other carbonates with liberation of carbonic acid gas and forma- tion of chlorides. Physical Properties. Hydrochloric acid can be con- densed to a liquid under considerable pressure and in this form has a specific gravity of 1.27. A liter of the gas, 76 GENERAL CHEMISTRY. under standard conditions, weighs 1.643 grams. As men- tioned, the gas is extremely soluble in water, it being pos- sible to prepare at a low temperature a solution which contains 45 per cent by weight of the acid. This solution is not stable at higher temperatures. A solution of 42 per cent strength can be made to keep. Our strongest com- mercial acid has usually a strength of about 40 per cent. The stronger grades fume when exposed to the air because of a combination of the real acid with the moisture present. Uses. Hydrochloric acid is employed for numerous purposes in the chemical laboratory and in technical oper- ations on the large scale. Much of the crude by-product, mentioned above, is employed in the manufacture of bleaching powder. ELEMENTS AND COMPOUNDS. Enough work has been done thus far by the student to make him acquainted with certain fundamental differ- ences between bodies. Oxygen has been prepared by sev- eral methods, and it was found that it could be readily combined with several other substances. Hydrogen and chlorine likewise, were secured in the free pure condition and in turn were united with other bodies to form new substances. Nowhere has anything been said about the decomposition or breaking up of oxygen, hydrogen and chlorine themselves. The question might naturally occur to the student, why, in experimenting with these three bodies, has no experiment been given in which they in turn should be decomposed. The fact is that up to the present time no means have been found by which these three substances, and many others to be mentioned later, can be resolved into any- thing simpler. The numerous and powerful methods of decomposition known to chemists have been applied in vain to the splitting of these bodies, and hence they have come to be regarded as the real elements of the material world. They cannot be decomposed, it appears, but they can combine to form other substances. The new sub- GENERAL CHEMISTRY. 77 stances are called compounds, and of these we have al- ready had in our work many illustrations. Water, common salt, hydrochloric acid, sulphuric acid, mercuric oxide, po- tassium chlorate, carbon dioxide, hydrogen dioxide and other bodies produced or used were shown by experiment to be compound in their nature. In all these cases the existence of at least two elements was shown or could be inferred with certainty from the experimental results. There appear to be about seventy-four of these ele- mentary substances, while the number of compound bodies known is enormously large. That hydrogen, oxygen, chlorine and the other so-called elements are really ele- mentary, that they can never be decomposed, yielding other substances, we cannot safely affirm ; indeed, water and many other compound bodies were once looked upon as elements. But this much may be safely said, that with the means now at our command, we cannot decompose them, and, therefore, for all practical purposes they must be looked upon as elementary. ATOMS AND MOLECULES. A systematic or scientific study of chemical phenomena began toward the end of the last century, and, as already pointed out, the conditions under which many of the ele- ments combine were soon recognized. It was found, among other things, that the power of combination is limited; in other words that the elements can be made to unite, as a rule, in certain proportions only. Several cases were known in which metals combined with oxygen in more than one proportion, but even the crude analyses of the time were sufficient to show that in the more highly oxidized bodies the amount of oxygen present is a multi- ple of that in the lower. In the two oxides of nitrogen known at the beginning of this century, it was found that one contains just twice as much oxygen as the second. The same relation was pointed out for the two compounds of carbon and oxygen known, and John Dalton found that in two compounds of carbon and hydrogen known the ratio of the weight of the carbon to that of the hydrogen is just 78 GENERAL CHEMISTRY. twice as great in one case as in the other. These observed facts naturally caused much speculation among chemists, butDalton was the first to propose a satisfactory hypoth- esis to account for them. From the earliest times philosophers were familiar with the idea that matter exists ultimately in the form of minute indivisible particles called atoms, and although this view was not regarded in general as of fruitful importance, Dalton was able to develop it further and make finally much of it. After applying the conception to the explana- tion of several purely physical phenomena, he employed it to account for the formation of chemical compounds by the union of minute particles or atoms of constant weight. According to him an atom of hydrogen unites with an atom of oxygen to form water; an atom of hydrogen with an atom of nitrogen to form ammonia; an atom of hydrogen with an atom of carbon to form ethylene. A contemporary of Dalton, Thomson, explaining the views of the latter, mses this language: " One atom of a body, a, unites with one atom of a body, b, or with two atoms of it, or with three, four, etc., atoms of it. The union of one atom of a with one of b produces one compound, the union of one atom of a with two atoms of b produces another compound, and so on." "We have no means of demonstrating the number of atoms which unite together in this manner in every com- pound; we must, therefore, have recourse to conjecture. If two bodies unite only in one proportion, it is reasonable to conclude that they unite atom to atom. Hence it is most likely that water is composed of one atom of oxygen and one atom of hydrogen; oxide of silver, of one atom of silver and one atom of oxygen; and oxide of zinc, of one atom of zinc and one atom of oxygen." "If we know the number of atoms of which a body is combined, and the proportion of the constituents, there is no difficulty in determining the proportional weight of the atoms of which it is composed. Thus, if water be com- posed of one atom of oxygen and one atom of hydrogen, and if the weight of the oxygen in water is to that of the hydrogen as 7^ to 1, then it follows that the weight of an GENERAL CHEMISTRY. 79 atom of oxygen is to that of an atom of hydrogen as 7^ to 1." The above quotations express clearly Dalton's notion of the combination of atoms to form larger groups, which we now call molecules. Because of lack of sufficient ex- perimental data he was led to assume that compounds formed by the union of atoms are in many cases simpler than we now'have reason to consider them, but this is a detail which does not detract from the theory. Dalton recog- nized that the weights of these minute atoms must be exceedingly small and beyond the reach of practical deter- mination. He therefore proposed a new system of weights, the weight of the atom of hydrogen being taken as the standard and called unity. The weights belonging to this system are commonly called the atomic weights, and one method of arriving at their value is suggested in the above quotation; other methods will be pointed out later. Since the time of Dalton many new elements have been discovered and exact analytical methods have been per- fected by which the weights of their ultimate atoms on the hydrogen scale may be readily found. For certain prac- tical reasons Berzelius, a contemporary of Dalton, sug- gested the atom of oxygen as the standard and proposed to call its weight arbitrarily 100. This suggestion did not meet with general favor. To-day, however, many chem- ists agree with Berzelius in his reasons for preferring the atom of oxygen, rather than the atom of hydrogen, as the standard, but place its weight at 16, which is very nearly its true weight on the hydrogen scale. The table given below contains a list of the elements with their atomic weights on the hydrogen scale and on the oxygen scale. This table has been calculated by Prof. F. W. Clarke, of the U. S. Geological Survey, and embraces the results of the latest and most accurate determinations. In the fifth col- umn of the table are given some approximate values ob- tained by rounding off the numbers of the fourth column. These approximate values are convenient and sufficiently accurate for the calculation of problems to be given later, and also for the illustrations which follow. The first col- umn contains the name of the element and the second col- umn the symbol by which it is represented. 80 GENERAL CHEMISTRY. Table of Atomic Weights. Name. Symbol. H=l O=16 Approx. Aluminum Al 26 91 27 11 27 1 Antimony Sb 119.52 120.43 120 4 Argon ....... Ar 40 Arsenic ... As 74 44 75 01 75 Barium Ba 136 39 137.43 137 4 Beryllium Be 9 01 9 08 Bismuth , Bi 206.54 208.11 208 1 Boron B 10.86 10.95 11.0 Br 79.34 79.95 80 Cd 111.10 111.95 111.9 CsGsium Cs 131 89 132 89 Calcium Ca 39.76 40.07 40.1 Carbon c 11 92 12 01 12 Cerium . ... Ce 139.1 140 2 Chlorine Cl 35 18 35 45 35 5 Chromium . . .... Cr 51.74 52 14 52 1 Cobalt Co 58.49 58 93 58 9 Columbium Cb 93.02 93.73 CoDDer Cu 63 12 63 60 63 6 Erbium *i Er 165.06 166.32 Fluorine F 18 91 19 06 19 Gadolinium Gd 155 57 156 76 Gallium Ga 69.38 69 91 Germanium Ge 71.93 72.48 Gold Au 195.74 197.24 197 2 Helium He 4 Hydrogen H 1.00 1.008 1.0 Indium In 112 99 113 85 Iodine I 125.89 126.85 126.9 Indium Ir 191.66 193 12 Iron . Fe 55 60 56 02 56 Lanthanum La 137.59 138 64 Lead Pb 205.36 206 92 206 9 Lithium Li 6.97 7.03 7 Magnesium . Me 24 10 24 28 24 3 Manganese Mn 54 57 54 99 55 Mercury Hg 198.49 200 00 200 Molybdenum Mo 95.26 95.98 96 Neodymium Nd 139.70 140.80 Nickel Ni 58.24 58.69 58.7 Nitrogen ... N 13 93 14 04 14 Osmium . . Os 189.55 190 99 GENERAL CHEMISTRY. Table of Atomic Weights. Continued. 81 Name. Symbol. H=rl O=16 Approx. O 15.88 16.00 16.0 Pd 105.56 106.36 Phosphorus P 30.79 31.02 31.0 Platinum Pt 193.41 194.89 194.9 K 38.82 39.11 39.1 Praseodymium . , .. Pr 142.5 143 6 Rhodium . . Rh 102 23 103 01 Rubidium Rb 84 78 85 43 Ruthenium Ru 100 91 101 68 Samarium Sm 149 13 150 26 Scandium . ... Sc 43 78 44.12 Selenium ... Se 78.42 79 02 79 Silicon Si 28 18 28 4 28 4 Silver Ag 107 11 107.92 107 9 Sodium . Na 22 88 23.05 23 Strontium . . Sr 86 95 87 61 87.6 Sulphur o s 31.83 32 07 32 1 Tantalum Ta 181 45 182.84 Tellurium Te 126 52 127.49 127 5 Terbium . . Tb 158 8 160.0 Thallium ... Tl 202.60 204 15 Thorium Th 230 87 282 63 Thulium Tm 169 4 170.7 Tin Sn 118.15 119 05 119 Titanium Ti 47.79 48.15 Tungsten W 183.43 184.83 Uranium u 237 . 77 239.59 239 . 6 Vanadium. V 50.99 51.38 Ytterbium Yb 171.88 173.19 Yttrium Yt 88 35 89 02 Zinc Zn 64 91 65 41 65 4 Zirconium Zr 89.72 90.40 Use of Symbols. We have reached a point now in our work where very great help is derived from the use of symbols representing the substances dealt with, and the student is advised to learn those of the important elements in the table. The first use of a symbol is as an abbreviation of the name of a substance dealt with. Thus, we use H as standing for 82 GENERAL CHEMISTRY, or representing hydrogen, O as representing oxygen, S as representing sulphur, and so on. Employed in this man- ner, we use the symbol merely to save time or space in writing. But there is a second and much more important application of these letters as representing something. H stands for the smallest weight of the element, hydrogen, which can exist in any compound or take part in any reac- tion, O for the smallest weight of oxygen, S for the small- est weight of sulphur combining or existing in the same manner. According to the definition given above, these symbols, therefore, represent weights of the several sub- stances corresponding to the atomic weights. In all of our calculations O stands for 16 parts of oxygen, H for 1 part of hydrogen, Cl for 35.5 parts of chlorine, Na for 23 parts of sodium, and so on. Molecules and Molecular Weight. The atoms of the elements mentioned above combine among themselves to form groups called molecules. Two atoms of hydrogen unite with one atom of oxygen to form a molecule of water; an atom of oxygen forms with an atom of mercury a molecule of mercuric oxide; two atoms of carbon, six atoms of hydrogen and one atom of oxygen in combina- tion constitute a molecule of alcohol. We represent mole- cules by uniting the symbols of their component atoms. Thus, for the above illustrations: H 2 O, HgO, C 2 H 6 O. We call this combination of symbols a formula. A symbol is, therefore, arbitrarily taken to represent an atom, while a formula represents a molecule. We may now apply some of these facts in explanation of experiments in the preceding chapters. In our experi- ment on the decomposition of mercuric oxide, with libera- tion of oxygen, we may represent what takes place by this equation: HgO := Hg + O M o r xide C =Mercury-f Oxygen. This equation tells us that 216 parts of the compound, mercuric oxide, yield when heated 200 parts of the ele- GENERAL CHEMISTRY. 83 ment, mercury, and 16 parts of the element, oxygen. In the decomposition of potassium chlorate we have a more complex case. This is a combination of potassium, oxygen and chlorine from which the oxygen may be readily sep- arated by heat. A study of the compound, potassium chlorate, shows that it contains its elements in these pro- portions by weight: oxygen, 48 parts; chlorine, 35.5 parts ; potassium, 39.1 parts. The atomic weights or combining weights are represented here, three times for oxygen, once for chlorine and once for potassium. We therefore write as the formula of our compound, KC1O 3 . When we de- compose this we find that all of the oxygen is given off and that we have a solid substance left which contains potassium and chlorine in the proportions in which they were found in the original compound. We therefore write this equation, as expressing the results of our experiment : KC1O 3 = KC1 -f O 3 Potassium _ Potassium I Qxvgon chlorate ~ chloride The student must early recognize this fact, that a chem- ical equation is always written to show, in compact form, what experiment proves has taken place or must take place under proper conditions. Chemical symbols cannot be combined at random, we cannot perform operations on them as we do with algebraic symbols, but when we write them on the left hand side of our equality sign we simply name the substance or substances on which some experi- ment is to be performed. After the experiment we are able to complete the equation, and then we write down on the other side of the equality sign what has taken place. The above equation shows correctly, only tne relations by weight between the substance taken and the products. Careful experiments have made it plain that the reaction really takes place in two stages, as represented by the fol- lowing equations : 2KClO 3 ^KClO 4 -fKCl+O 2 . 84 GENERAL CHEMISTRY. Potassium chlorate yields at first a substance known as potassium perchlorate with potassium chloride and a relatively small amount of oxygen, in fact just one-third, by weight, of that in the original compound. At a high temperature the potassium perchlorate, represented by the formula KC1O 4 , breaks up into more potassium chloride and oxygen, thus, KC10 4 :=KC1+0 4 . Combining the two equations we can therefore write: 2KClO 3 = KCl-fKCl+O 2 +O 4 , or better, 2KC1O 3 =2KC1+3O 2 . Why we write 3O 2 instead of O 6 is a question which cannot be satisfactorily answered at this point, but will be taken up later in consideration of other experimental re- sults; but another simple matter in connection with the method of writing equations must be explained here. The student observes that numerals are employed in two posi- tions in these equations. Large figures are written before the formulas of compounds, while small figures in several cases seem to be to the right and a little below certain symbols. This is a purely conventional arrangement, and the meaning conveyed could be just as well expressed in some other manner. It has been agreed by chemists to consider the large numerals as multiplying the whole com- pound which follows, while the small figures are taken as referring only to the symbol of the element immediately preceding. Thus, in 2KC1O 3 we have twice the whole group, while the small 3 indicates that we have in each group the combining weight of oxygen taken three times. The effect of the large numerals, however, is not carried beyond a sign of addition or subtraction. In the equation, the large 2 at the beginning refers only to the KNO 3 , and not to H 2 SO 4 . GENERAL CHEMISTRY. 85 We may now express, by the use of symbols, the reac- tions which took place in some of the experiments on oxy- gen. We have C + 2 = C0 2 . Carbon + Oxygen = Carbon S + O 2 = SO 2 . Sulphur + Oxygen^ P 4 Phosphorus = 2P 2 6 . In the first case we have the union of 12 parts of car- bon with 32 parts of oxygen. In the second case 32.1 parts of sulphur combine with 32 parts of oxygen, while in the third case 124 (4X31) parts of phosphorus combine with 80 (5 X 16) parts of oxygen. The combining weight of zinc has been found to be 65.4, and we find that 65.4 parts of zinc act on 98.1 parts of sulphuric acid, liberating 2 parts of hydrogen. We there- fore write the equation, Zn -fH 2 S0 4 = ZnS0 4 + H 2 . Zinc +>Sf C = su ^ h C ate + H y drogen. This equation, like all the others, is intended to express the results of experiments. We find that sulphuric acid is composed of hydrogen, oxygen and sulphur in certain pro- portions shown by the formula H 2 SO 4 , and that at the end of the experiment we have in solution a compound which contains, in place of hydrogen, zinc combined with oxygen and sulphur, in the proportions shown by the formula, ZnSO 4 , and which is known as zinc sulphate. By the sym- bol H we represent 1 combining weight of hydrogen, and by H 2 we represent two such weights. By our equation, therefore, we express this fact, that in the solution of 65.4 parts of zinc we liberate 2 parts by weight of hydrogen. To generate 2 grams of hydrogen gas by this method we must dissolve 65.4 grams of the metal zinc. 86 GENERAL CHEMISTRY. It has been already mentioned that we may use iron in- stead of zinc for the generation of the hydrogen gas. In this case we find that 56 parts of iron produce the same amount of hydrogen that we obtain from the 65.4 parts of zinc. The 56 represents, in fact, the combining weight of iron, and we may write as expressing the last reaction, Fe +H 2 SO 4 = FeSO 4 -f H 2 Tr/,, I Sulphuric Ferrous ( j Iron + acid = sulphate + Hydrogen. 56 + 98.1 = 152.1 -f- 2. The sum of the weights on the left hand side of the equation is equal to the sum on the right. There is neither a gain nor loss of matter, but merely a rearrangement of elements in the compounds. It appears from these illus- trations that zinc and iron have the power of displacing the hydrogen in the acid used. Had we used hydrochloric acid instead of sulphuric exactly the same behavior would have been observed. Hydrogen would have been displaced in quite the same manner, and 2 parts by weight for 65.4 parts of zinc, or 56 parts of iron dissolved. In general, it may be said that zinc and iron displace hydrogen in many bodies called acids, and always in the proportions given. We come now to a consideration of the reactions by which chlorine and hydrochloric acid were produced, and here again we deal with the results of exact experiments. It has been found that sulphuric acid is able to decompose sodium chloride, or common salt, in a manner illustrated by the following equation: Na 2 SO 4 -h 2HC1 Sulphuric , Sodium ___ Sodium i Hydrochloric acid ~t~ chloride ~ " sulphate "1 acid 98.1 + 117 = 142.1 -f 73. Experiment actually shows us that for an amount of sulphuric acid represented by the sum of the combining weights of its elements, that is for 98.1 parts by weight, GENERAL CHEMISTRY. 87 we require 117 parts of salt, that is, twice the sum of the combining weights of the sodium and chlorine. When we employ the pure materials in exactly these proportions, and aid their action on each other by heat, we find at the end of our experiment that we have neither sulphuric acid nor salt, but two new substances, one of which is the hydrochloric acid, which we collect in water, and the other a white solid substance which remains in the decomposing vessel, and which we call sodium sulphate. There are no other products in the reaction. We call this a reaction of double decomposition, inasmuch as we start with two compound bodies which react on each other to form two new compound bodies. We can illustrate this double decomposition by a diagram, as follows: C- -D C ^ "ND Before the reaction, one compound bod} 7 is made up of the parts A and B, and the other compound body of the parts C and D. But after the decomposition we have a new compound body, with A and D as its parts, and another with C and B. The reaction between salt and sulphuric acid is a typical one of double decomposition and well illustrates many which are to follow. From our hydrochloric acid, as made above, we sepa- rate the chlorine by another process described. This is somewhat more complex, but its exact nature may be read- ily illustrated as follows: MnO 2 + 4HC1* s MnCl 2 +2H 8 O+ C1 2 Manganese _1_ Hydrochloric Manganese_|_ W iff>r 4-rhlnrinp dioxide acid ~ chloride ' 87 -f 146 = 126 -f 36 -f 71. The manganese dioxide in the above is made up of the combining weight of manganese plus twice that of oxygen. Now experiment shows that to decompose this completely 88 GENERAL CHEMISTRY. we need of hydrochloric acid four times the sum of the combining weights of hydrogen and chlorine. Less would not be sufficient to complete the decomposition of the manganese dioxide. The experiment may be performed in such a manner as to show that water is liberated, and exactly how much. For the amount of manganese dioxide assumed to be taken we find of water just twice the sum of the combining weights of hydrogen and oxygen, the weights taken or obtained being expressed in any con- venient standard, as in grams. The chlorine liberated is just half of that contained in the original hydrochloric acid used, which fact is expressed also in the equation. When chlorine is liberated by the action of salt, sul- phuric acid and manganese dioxide on each other, the whole of that element in the salt may be obtained. The reaction takes place in two stages possibly, the first involv- ing the formation of hydrochloric acid and the second its decomposition. The following equation shows the quanti- tative relations existing between the compounds taken and obtained. + 2NaCl+2H 2 S0 4 = MnS0 4 +Na 2 S0 4 + 2 87 + 11? + 196.2 = 151.1 + 142-1 + 36 +71. From the above it appears that the manganese dioxide used, and chlorine obtained stand to each other in the re- lation of 87 to 71. As both weights are referred to the same basis or standard, the proportion must hold good if we refer them to a new standard. If we take the gram as our unit it is true that 87 grams of manganese dioxide must be used for the liberation of 71 grams of chlorine, and if we weigh in pounds ortons.the same ratio must still exist. It is easy, therefore, to tell how much manganese dioxide must be used to liberate any given quantity, as 100 grams of chlorine. It is evident that the correct answer must be given by the following proportion : 87 : 71 :: x : 100. x = 122.5 grams. GENERAL CHEMISTRY. 89 To find the amount of salt required we make another proportion : 117 : 71 :: x : 100. x=164.8 grams. A third proportion shows the amount of sulphuric acid necessary for the formation of the 100 grams of chlorine : 196.2 : 71 :: x : 100. x = 276.3 grams. Aided by the above explanations the student should now be able to understand what follows. From this point on all important reactions will be represented by equations, and these should be thoroughly studied. The student should keep in mind, however, that equations are not drawn from the imagination, but represent, properly, the results of experiments. He should practice writing them as an aid to memorizing important reactions, and especially because of their value in the solution of even the simplest chemical problems, as illustrated by the examples given above. In the present chapter no attempt will be made to explain methods by which the atomic weights are found. In one to follow, however, after the student has become more familiar with chemical facts, something on this topic will be given. CHAPTER IV. COMPOUNDS OF CHLORINE WITH OXYGEN.- BRO- MINE, IODINE, FLUORINE AND THEIR COMPOUNDS. TN the last chapter the element, chlorine, and its com- pound with hydrogen have been described. In this chapter a few other important combinations must be referred to. OXIDES AND ACIDS OF CHLORINE. Three compounds of chlorine with oxygen are known, but they cannot be formed by direct union. Chlorine Monoxide and Hypochlorites. The first one is called chlorine monoxide, and is represented by the formula C1 2 O. It is a yellowish brown gas with an odor suggesting chlorine, and may be made by passing dry chlorine over the red oxide of mercury, freshly pre- cipitated and dried, and contained in a glass tube. This equation expresses the combination : The gas can be easily condensed to a liquid, but this is not stable. It is very soluble in water, forming a new acid, called hypochlorous acid : C1 2 O + H 2 O = 2HOC1. The acid solution is not stable; if the gas is led into an GENERAL CHEMISTRY. 91 alkali solution, however, an important body called a hypo- chlorite is formed: C1 3 O+2KOH = 2KOC1-|-H 2 O. Some of the hypochlorites are well known and useful substances. Calcium hypochlorite is the active constitu- ent of bleaching powder. Sodium and potassium hypochlo- rites are used in the laboratory and in medicine. These hypochlorites are easily decomposed by hydrochloric or sulphuric acid with liberation of chlorine. Practical appli- cation is made of this in bleaching by bleaching powder, which is illustrated by the following experiment: Ex. 72. Pour some dilute sulphuric acid over a few grams of bleaching powder in the bottom of a large beaker which then cover with a piece of glass or a card. Observe that greenish yellow fumes soon collect in the beaker. Moisten now a piece of bright calico, as already described, and hang it in the beaker of gas. The calico will fade as before. The manufacture of bleaching powder will be referred to later. It is made by passing chlorine gas over slaked lime, and is essentially a mixture of calcium hypochlorite, CaO 2 Cl 2 , and calcium chloride, CaCl 2 , in about equal proportions. Chlorine Dioxide. This is a heavy dark yellow gas usually made by the decomposition of potassium chlorate by sulphuric acid. Chloric acid, HC1O 3 , is formed first and this decomposes on slight warming. The gas can be condensed to a liquid at a low tempera- ture. It is not stable, often decomposing with explosive violence. It dissolves rather readily in water but does not form a new acid. Chlorine Trioxide. This is a greenish yellow gas having the composition, C1 2 O 3 , and is made by several proc- esses depending on the reduction of chloric acid, HC1O 3 . 92 GENERAL CHEMISTRY. It is decomposed by warm water forming a mixture of hydrochloric and chloric acids, and with cold water yields chlorous acid, as below. Chlorous Acid, HC1O 2 , is not known in the pure state, but certain salts, called chlorites, are known and these correspond to the acid. Solutions of chlorous acid result when the trioxide is dissolved in water: C1 2 O 3 +H 2 O = 2HC10 2 . Chloric Acid. This is the best known of the oxygen acids of chlorine and may be prepared by decomposing barium chlorate by means of sulphuric acid : Ba(ClO 3 ) 2 + H 2 SO 4 =:2HClO 3 -hBaSO 4 . As barium sulphate, BaSO 4 , is a very insoluble precip- itate it is easy to obtain a pure solution of the chloric acid by pouring off the supernatant liquid. The con- centrated acid is a strong oxidizing agent, has a pungent odor and decomposes readily when heated, yielding a new acid known as perchloric acid, HC1O 4 , along with oxygen and chlorine. Corresponding to chloric acid, we have the well-known salts called chlorates, of which potassium chlorate, KC1O 3 , is the best illustration. The chlorates are all soluble in water and decompose when heated, yielding oxygen. The decomposition takes place in two stages, however; in the first perchlorate is formed: 2KC1O 3 =KC10 4 In the second stage of the eaction the perchlorate is decomposed, yielding more oxygen and chloride. Perchloric Acid. As potassium perchlorate is but slightly soluble in water, advantage is taken of the above reaction in preparing perchloric acid, HC1O 4 . When the GENERAL CHEMISTRY. 93 chlorate is heated until the evolution of oxygen begins the first stage of the reaction may be considered as completed, practically. If the mass is now cooled, powdered and extracted with water the perchlorate is left while the chloride goes into solution. This perchlorate distilled with strong sulphuric acid yields perchloric acid, which is a heavy, volatile liquid having a great affinity for water. It is a powerful oxidizing agent and decomposes immediately when brought in contact with most organic substances. The perchlorates are all soluble in water and they differ from the chlorates in not being decomposed by hydro- chloric acid. It appears from the foregoing that we have four chlorine acids containing oxygen. The names and formulas of these are : Hypochlorous acid, HC1O, Chlorous acid, HC1O 2 , Chloric acid, HC1O 3 , Perchloric acid, HC1O 4 . It will be observed that the names differ through certain prefixes and terminations and it will be seen later that the same are used in the designations of all other acids. OTHER CHLORINE COMPOUNDS. Chlorine combines indirectly with nitrogen to form a very explosive substance known as nitrogen chloride. It forms a number of important combinations with carbon and with carbon and hydrogen, to be mentioned later. A very important compound with oxygen and nitrogen will be described in the next chapter. BROMINE. Occurrence. Bromine is an important element which never occurs free in nature. It is found in several bro- mides in spring waters, and to a slight extent in sea water. 94 GENERAL CHEMISTRY. History. Bromine was discovered in 1826 by Balard in the mother liquor left after crystallization of salt from evaporated sea water. The discoverer was able to show the important analogies existing between this element and chlorine and iodine. Preparation. Much of our bromine is obtained from the residues left on crystallizing salt from concentrated brine of certain salt springs. The bromine is left in these mother liquors in the form of bromides, which are more soluble than the common salt, and may be liberated by several reactions, of which two illustrations will be given. Large quantities of bromine are produced at the Michigan salt wells and also from the salt deposits of Stassfurt, Germany. Ex. 73. Dissolve a small crystal of sodium or potassium bromide in water in a test-tube, and add gradually, a drop at a time, some chlorine water. Use for this purpose the chlorine water saved from a former experiment. When the first drop of chlorine water mixes with the solution of bromide it produces a reddish-yellow color which deepens to red as more of the reagent is added. If the bromide solution is weak and the chlorine water strong, the red color will finally disap- pear by continued addition of the latter. The chlorine water decom- poses the bromide, liberating bromine. In the above experiment we have illustrations of sev- eral important points. First, of the liberation of free bro- mine. Potassium bromide is a combination of bromine with potassium, which we represent by the formula KBr. Sodium bromide is represented by NaBr. Assuming that we are dealing with the former we express the whole reaction by this equation: KBr -f Cl = KC1 + Br 119.1 + 35 -5 = V4.60 -f- 80. The equation shows just what an exact quantitative ex- periment would have revealed to us, viz., that 35.5 parts by weight of chlorine are required to completely decompose 119.1 parts of potassium bromide with liberation of 80 parts of bromine. The weight of chlorine taken and that of bromine obtained are chemically equivalent, but, as the GENERAL CHEMISTRY. 95 result shows, the chlorine is able to displace the bromine. We are no more able to give an exact reason for this dis- placement than we are to assign a reason for other chem- ical decompositions already illustrated. But we are accus- tomed to say that the chlorine has a greater affinity for the potassium than the bromine has and is therefore able to drive it out from its combination. Of the real nature of this chemical affinity we know but little. The above experiment illustrates the marked activity of chlorine in another manner. It was shown that a great excess of the chlorine water discharged the color of the free bromine. This loss of color is due to two causes. First, to the combination of the excess of chlorine with the bromine liberated, forming bromine chloride, and second, to the oxidation of some of the bromine to bromic acid, in presence of water, which is illustrated by this equation: Br+3H 2 O + 5Cl = 5 HCl+HBrO 3 . Hydrochloric and bromic acids result. Bromine can be readily liberated from bromides by a reaction analogous to that employed for the preparation of chlorine from chlorides, that is, by the use of sulphuric acid and manganese dioxide. The following experiment will illustrate this: Ex. 74. In a 300 Cc. flask mix about 2 Gm. of powdered potassium bromide with 4 or 5 Gm. of commercial powdered manganese dioxide. Add a little water and shake until the mixture becomes uniformly dis- tributed. Then add 50 Cc. of dilute sulphuric acid and close the flask with a stopper through which passes a long delivery tube bent down to dip into a small flask or beaker of cold water. The flask with the above described mixture must stand on a sand-bath or wire gauze, which is then heated by a lamp. Red vapors are generated in the flask which distill over and dissolve in the water in the small receiving flask or beaker. Continue the application of heat as long as these red vapors are evolved. Then remove the receiving flask and withdraw the lamp from the other flask. The reaction which takes place here is illustrated by this equation: 2KBr+MnO 2 4-2H 2 SO 4 =: Br 2 4-MnSO 4 -f-K 2 SO 4 +2H 2 O. 96 GENERAL CHEMISTRY. This is seen to be similar to the chlorine reaction. A very large excess of manganese dioxide is taken in the ex- periment, in order to secure the complete decomposition of the bromide without liberation of hj'drobromic acid. On the large scale this process is applied to the manu- facture of bromine from the mother liquors of salt works. As these liquors contain much chloride, chlorine is first liberated and this serves to free the bromine. Some bromine chloride is always produced in the operation, but this is more volatile than the bromine and can be sepa- rated by distillation. In the experiment just described, bromine was collected in water. It is somewhat soluble in water, as shown by the fact that at first all that distilled over went into solu- tion. Before the end of the experiment, however, unless too much water was taken a part of the bromine settles out as a dark red drop. Use this aqueous solution of bro- mine for tests as follows: Ex. 75. Bromine bleaches as does chlorine, but with less activity. Test this by use of colored calico, and also with solutions of organic col- oring matters, litmus and cochineal, for instance. Bromine is readily soluble in chloroform, carbon disul- phide, ether and other liquids, and may be withdrawn from aqueous solution by them. Ex. 76. Pour some of the bromine water, made above, into a test-tube, and add about one-tenth its volume of carbon disulphide. Close the tube with a cork and shake thoroughly. On standing, the disulphide speedily collects at the bottom of the tube, and it will be seen that it is highly colored by the absorbed bromine, while the water above is much lighter colored than before, or it may be even colorless. This beautiful reaction is employed in the detection of small traces of bromides in spring water, the bromine being first liberated by means of a small amount of chlorine. Physical Properties. Bromine boils at 63, and freezes about 7. At it has a specific gravity of 3. 18. At 15 it dissolves in 33 parts of water. Uses. Bromine is employed in the preparation of bromides, several of which are used in medicine. It is GENERAL CHEMISTRY. 9? used also in making certain reagents employed in labora- tories and in making a number of valuable organic prepa- rations. BROMINE AND HYDROGEN. Under certain conditions these two elements may be directly united, but not as readily as is the case with chlo- rine and hydrogen. Hydrobromic acid, HBr, results. This acid cannot be made in pure condition by the reaction employed in the manufacture of hydrochloric acid, that is, by the decomposition of a bromide by means of strong sul- phuric acid, according to the following equation : The hydrobromic acid as liberated is partially decom- posed by the excess of strong sulphuric acid, free bromine and sulphurous oxide being formed. This can be illus- trated as follows : Ex. 77. Take some small crystals of potassium bromide in a test- tube, and pour over them a little strong sulphuric acid. An escape of gas is seen to follow, which has a yellowish color, due to fi'ee bromine, the hydrobromic acid itself being a colorless gas. Pure hydrobromic acid may be made, however, by using phosphoric acid instead of sulphuric acid, and by a reaction illustrated by this equation : PBr 3 +3H 2 O = 3HBr-{-H 3 PO 3 . Phosphorous bromide is decomposed by water, yielding hydrobromic acid and phosphorous acid. Instead of using pure phosphorous bromide it is customary to add bromine very slowly to a mixture of red phosphorus and water in a suitable apparatus arranged in such a manner that the gas as it escapes may be absorbed in water. Properties. In most of its important properties hy- drobromic acid resembles hydrochloric acid. It is a gas, and very soluble in water. One cubic centimeter weighs .003645 Gm. at and normal pressure. It is readily de- 98 GENERAL CHEMISTRY. composed by chlorine and it unites with alkalies forming bromides. Hydrobromic acid and all the soluble bromides give a precipitate when treated with a solution o'f silver nitrate as shown below : Ex. 78. Prepare a dilute solution of potassium bromide by dis- solving a small crystal in water. Add to this solution a few cubic cen- timeters of a dilute solution of silver nitrate. A yellowish white, curdy precipitate forms which soon settles to the bottom of the vessel in which it was produced. This precipitate is not soluble in nitric acid, and to a limited extent only, in dilute ammonia. The compounds of bromine with hydrogen and oxygen are not of sufficient importance to be taken up in this place in detail. No oxides are known, but two oxygen acids, hypobromous acid, HBrO, and bromic acid,HBrO 3 , are known. Some of the hypobromites, the salts formed from HBrO, are used as reagents. IODINE. Iodine is a very important element resembling chlorine and bromine in certain chemical properties, but is a steel gray solid at the ordinary temperature. It is far less abundant than either chlorine or bromine in nature, occur- ring in some springs, but mainly in sea water, from which it is taken up by certain seaweeds. The iodine of com- merce is largely obtained from the ash produced by burn- ing these weeds. It occurs also in small amount in com- pounds called iodates, which occur with Chili saltpeter. History. Iodine was discovered by Courtois, a French chemist, in 1812. It was found in the mother liquors left after the extraction of sodium salts from kelp or the ash of seaweeds. Until recently this kelp, or varec, was the source from which practically all iodine was obtained. A larger proportion is now produced from the mother liquors occurring in the refining of Chili saltpeter. As produced from seaweed iodine is obtained mainly from the coasts of Scotland and northern France. GENERAL CHEMISTRY. 99 Preparation. As found in the ash from seaweed, the iodine occurs in the form of an iodide and can be sep- arated just as bromine is from bromides. The following experiments will illustrate this : Ex. 79. Dissolve a small crystal of potassium iodide in about five Cc. of water in a test-tube. Then add chlorine water, a drop at a time, which produces a brown color, and finally, if the solution is not too weak, a precipitate of free iodine. An excess of chlorine water dis- charges the color as in the corresponding case with a bromide, and for the same general reasons. The decomposition is illustrated by this equation : 35.5 parts of chlorine replace 126.9 parts of iodine. Iodine is liberated, also, by the reaction with manga- nese dioxide and sulphuric acid, which can be easily illus- trated by a simple experiment as follows : Ex. 80. Mix about a gram of powdered potassium iodide with two or three times this weight of powdered manganese dioxide in a flask of 300 to 400 Cc. capacity. Pour in 5 Cc. of dilute sulphuric acid and heat the flask on a sand-bath. Decomposition of the iodide takes place and deep violet colored vapors fill the flask. The vapors condense, in .part, on the upper and cooler portions of the flask. The decomposition is illustrated by the equation, 2 . K 2 SO 4 -fMnSO 4 -f2H 2 O-fI On the large scale the reaction is so carried out that the iodine distills over from the decomposing retorts of earthenware or firebrick, and condenses in cold receivers. Properties. Commercial iodine occurs as a steel gray crystalline solid. In its power of combination with metals iodine is less active than chlorine or bromine. The important properties of the substance may be shown by simple experiments. 100 GENERAL CHEMISTRY. Ex. 81. Heat a small crystal of iodine in a test-tube. The iodine vaporizes quickly, so that the whole tube may be filled with the violet colored vapors. That these vapors are heavier than air may be shown by holding the tube in front of a sheet of white paper as a background, and then turning the tube so that the vapor may flow down and show against the paper. When the tube cools, add a little distilled water and shake thoroughly. Iodine is slightly soluble in water, which is shown by the yellow color imparted in this test. To the aqueous solution add a few drops of cold starch paste. A beautiful blue results, due to the combination of the starch with iodine. Ex. 82. Iodine is much more soluble in alcohol than in water. Powder a small crystal of iodine and transfer to a test-tube. Then add about 2 Cc. of alcohol and shake thoroughly. The iodine dissolves, producing a brown solution known as the tincture of iodine. A drop of this tincture added to a beaker of water containing a little starch paste produces a blue color. Ex. 83. Iodine dissolves very readily in an aqueous solution of potassium iodide, which can readily be shown by adding a small amount of powdered iodine to potassium iodide solution. A dark brown liquid results. This is known as Lugol's solution, or the "compound solution of iodine " of the pharmacopoeia, when made with certain definite quan- tities of iodine, potassium iodide and water. Show that the addition of a great excess of water produces a precipitate in this solution, and that it gives the blue color with starch. Ex. 84. Dissolve minute crystals of iodine in carbon disulphide, ether and chloroform, and observe the colors of the solutions. Add chlorine water to a very dilute aqueous solution of potassium iodide in a test tube, until a brown color is formed, and then add several large drops of carbon disulphide and shake. The iodine is taken from the water by the disulphide, imparting to the latter a characteristic color. Pure iodine has a specific gravity of 4.95. It melts at about 115 and boils above 200. Uses. Iodine is employed in the preparation of iodo- form and several iodides used in medicine. It enters into the composition of many organic compounds. IODINE AND HYDROGEN. Hydriodic acid, HI, is a well-known and important substance, best made by the action of phosphorus and water on iodine by a process analogous to that employed in making hydrobromic acid. When potassium iodide is GENERAL CHEMISTRY. 101 distilled with sulphuric acid, pure hydriodic acid is not obtained, as a decomposition of this by the excess of sul- phuric acid follows, with liberation of iodine. In its chemical behavior this acid closely resembles hydrochloric and hydrobromic acxdfy but/it is le^Ss stable. It decomposes, liberating iodjajs. In water it, is,ex,trernely soluble, yielding a heavy solution?, 1 , \ .*-> ; **<>.' \\* \ ,',' Two oxygen acids of iodine are known; one of these is called iodic acid and is represented by the formula HIO 3 . It is a white crystalline solid, soluble in water, and is best made by the oxidation of iodine by strong nitric acid. When strongly heated it decomposes, yielding solid iodine pentoxide, I 2 O 5 , and this in turn dissolves in water repro- ducing the acid. H 8 + I 8 6 =2HIO S . Several iodates are known; sodium iodate is found in Chili saltpeter. Finally, a more highly oxidized compound of iodine is known and this .is called periodic acid, HIO 4 . It is a colorless crystalline solid, very soluble in water. Several unimportant compounds with chlorine and bromine are known, and also a singular compound with nitrogen which will be described later. FLUORINE. This is a gaseous element of which but little is known in the free state. Occurrence. It is found in nature in two important mineral compounds. One of these is calcium fluoride, CaF 2 , and is called fluorspar. The other is a so-called double fluoride, containing sodium and aluminum, AlF 3 -f- 3NaF, called cryolite. Fluorine is found in bones and especially in the teeth in small amounts. History. Some combinations of fluorine have been known for many years, but all attempts to isolate the ele- ment failed until quite recently. Moissan succeeded a few 102 GENERAL CHEMISTRY. years ago in liberating it by the electrolysis of hydrofluoric acid in a platinum tube at a very low temperature. Properties; Fluorine is characterized by its remarka- bly strcng'affiniuies.r Jt c combines with nearly all elements except o;&y,gen,,. It attacks glass to combine with its silicon and corrodes :metais at-the ordinary temperature quickly. Because of these peculiarities chemists have found difficulty in studying it rather than in decomposing its compounds. At the low temperature of Moissan's experiments it may, however, be set free in platinum. It is a yellow gas which decomposes water instantly forming hydrofluoric acid and oxygen. Many substances burn with the gas as they would with oxygen. It has recently been liquefied. FLUORINE AND HYDROGEN. An important combination of these elements is known. This compound is known as hydrofluoric acid, HF, and is very soluble in water. The solution is now an article of commerce and is sold for several purposes. It is usually prepared by the action of sulphuric acid on calcium fluoride, a native mineral substance found in quantity in several localities. This reaction is analogous to that by which hydrochloric acid is made from sodium chloride by means of sulphuric acid, and may be illustrated by this equation : 2HF Calcium _|_ Sulphuric Calcium i Hydrofluoric fluoride ' acid sulphate ' acid. The fact that hydrofluoric acid attacks glass may be shown easily by experiment. Ex. 85. In a lead dish, having a diameter of 5 centimeters or more, make a pasty mixture of strong sulphuric acid and powdered cal- cium fluoride. Place the dish on a sand-bath and by means of splinters of wood support over it a square of glass, both surfaces of which have been covered with wax. In the center of one of the waxed surfaces scratch some letters or figures and expose this surface to the action of the fumes which arise from the dish when it is heated. The temperature GENERAL CHEMISTRY. 103 must not be allowed to get high enough to melt the wax. After fifteen or twenty minutes remove the glass and scrape off the layers of wax. It will be observed that at the exposed points the figures or letters have become fixed in the glass by its corrosion. Much of our chemical graduated ware is marked in this manner. Etching is frequently carried out by immersing the glass article, properly protected by wax, in an aqueous solution of hydrofluoric acid or in a mixture of dilute sul- phuric acid and powdered fluorspar. No heat is applied, but a longer time must be given to complete the work. This reaction is due to the affinity of the fluorine for an element of the glass, the silicon. Some of this element is dissolved out by the hydrofluoric acid, forming silicon fluoride, SiF 4 . The aqueous solution of the acid cannot be kept in glass or iron vessels. It is handled on the small scale in bottles of hard paraffin or gutta-percha, and in large quantities in barrels coated with paraffin. Some of the fluorides are becoming important articles of commerce. The native calcium fluoride is largely used as a flux in the smelting of iron ores. GENERALITIES. The four elements, fluorine, chlorine, bromine and iodine, constitute a natural group, in which a variation in properties is closely related to a variation in atomic weight. Fluorine, the lightest element, has the strongest affinity for hydrogen and all the metals, but it forms no combination with oxygen. Iodine, the heaviest of the group, forms a stable compound with oxygen, while its combinations with hydrogen and the metals are very easily decomposed. Fluorine decomposes water immediately, iodine not at all. The decomposition by chlorine is much slower than by fluorine, while by bromine it is extremely slow. Of the three well known elements in the group, bromine stands between the lighter chlorine and heavier iodine in all important properties. In the following table some of the most important rela- tions of the four elements just considered are pointed out in form suitable for easy comparison : 104 GENERAL CHEMISTRY, F. Cl. Br. I. Atomic weight. 19.06 38.12 35.45 70.90 1.33 Decomposes it in light. HC1 C1 2 O C1 8 S C1O 2 HOC1 HOC1O HOClOa HOClOs 79.95 159.90 3.18 Decomposes it very slowly. HBr Vone known. HOBr HOBrO 2 HOBrOg J26.85 253.70 4.95 Does not decompose. HI IiO, HOIO 2 HOIOs Molecular weight Liquid density Action on water. . . Decomposes it very readily. HF None known. Hydrogen compounds Oxygen compounds Oxygen acids. . NATURE OF ACIDS. In the foregoing pages the term acid has been fre- quently employed, and from the experiments made or sug- gested the general composition and properties of these bodies have been indicated. Hydrochloric acid contains chlorine and hydrogen, hydrobromic acid, bromine and hydrogen, hydriodic acid, iodine and hydrogen, hydroflu- oric acid, fluorine and hydrogen. That sulphuric acid also contains hydrogen is evident from many experiments. It will be shown in the next chapter that the very common and important nitric acid is also a hydrogen compound. In general, it may be said here, acids are bodies charac- terized by containing hydrogen, which may be readily re- placed by metals to form a group of compounds known as salts. When zinc is dissolved in hydrochloric acid hydro- gen escapes and a salt called zinc chloride is produced. When zinc is dissolved in the other acids, zinc bromide, zinc iodide, zinc sulphate, etc., are formed, hydrogen in all cases being set free. GENERAL CHEMISTRY. 105 Acids neutralize solutions of bodies known as alkalies and bases, forming, as before, salts. We had an illustra- tion of this in the experiment in which hydrochloric acid was mixed with the solution of caustic soda. On evapora- tion, common salt, or sodium chloride, was left. In a simi- lar manner a mixture of hydrobromic acid with caustic soda would yield sodium bromide. Acids, bases and salts comprise by far the larger num- ber of substances considered in inorganic chemistry. In a following chapter the relations of these bodies to each other and to certain allied substances will be pointed out. CHAPTER V. NITROGEN AND THE ATMOSPHERE. GAS PROBLEMS. NITROGEN is a gaseous element found in the uncom- bined state in the atmosphere, and in combination widely distributed through plant, animal and mineral sub- stances. In vegetable tissues it is found in all alkaloids and in all the so-called proteid compounds. It occurs in the proteids of the animal kingdom also and in many sub- stances produced by animals. In the mineral kingdom it is found mainly in the substances called nitrates, of which common saltpeter and Chili saltpeter are the best illustra- tions. r History. Rutherford, in 1772, was apparently the first to suggest by actual experiment the presence in the atmos- phere of a gas incapable of supporting life and combustion. Scheele later, about 1774, came to the conclusion that the atmosphere must consist of two distinct gases, but it remained for Lavoisier, in 1775, to make a clear statement of the nature of the two important gases in the atmosphere. The one not supporting life or combustion he called azote, while the name nitrogen was suggested later by Chaptal. Preparation. We can obtain nitrogen from the air by separating in some manner the other important element, oxygen, from it. As oxygen enters readily into combina- tion with many substances, while nitrogen is inert, this can easily be done. In illustration, the following experiment may be made : GENERAL CHEMISTRY. 107 Ex. 86. We may take advantage of the reaction between phos- phorus and oxygen to free the nitrogen from the latter element. To this end dry a very small piece of phosphorus, not larger than a pea, and enclose it in a little cylinder of wire gauze. Attach a piece of iron wire to this cylinder, as a handle, and bend it into a U shape with one limb longer than the other. The gauze cylinder is attached to the shorter limb. Now ignite the phosphorus, hold it over a vessel of water, and then depress a wide mouth bottle of about 300 to 400 Cc. capacity, over the burning substance, so that the mouth of the bottle dips beneath the surface of the water. A volume of air is thus confined and exposed to the action of the burning phosphorus. In a few seconds the combustion is complete, when it will be found that the level of the water in the bottle is above that in the vessel, this being the case because water must ascend to take the place of the oxygen, united with the phosphorus. After the disappearance of the fumes of phosphoric oxide, by solution, the remaining gas may be tested. Withdraw the wire gauze, leaving the mouth of the bottle still under water, then close the mouth by a glass plate and bring it into the upright position on the table. Test the gas in the bottle as oxygen was tested, using, however, a burning taper or splinter in place of one merely glowing. The flame will be extin- guished, showing the inert nature of the gas. At the beginning of the above experiment the heat of the combustion expanded the air in the bottle and drove part of it out before it was completely acted on by the phosphorus. It follows, therefore, that the gas volume left does not accurately represent the proportion of nitro- gen in the original air. It will be shown later that the nitrogen should amount to very nearly four-fifths of the whole. The nitrogen as obtained by the above process is never quite pure, but by more elaborate methods it may be secured from the air in practically pure condition. We may pro- duce it in a pure state by the decomposition of certain compcunds containing it, and one such method will be illustrated here, in which we use ammonium chloride: Ex. 87. Make a mixture of powdered potassium dichromate and powdered ammonium chloride, using four parts by weight of the former to one part of the latter. With this mixture half fill an iron gas pipe re- tort, about 20 Cm. long and 1.5 Cm. in internal diameter, the arrange- ment of which is shown in the next figure. The pipe is closed with a cork and delivery tube, which dips be- neath the surface of water in a trough. On applying heat to the retort its contents decompose with liberation of nitrogen gas, which passes through the delivery tube, and may be collected by displacement of water in the usual manner. Fill several bottles with the gas and test 108 GENERAL CHEMISTRY. as follows: Remove the bottles with glass plates as in other cases. Into one bottle thrust a burning taper or piece of wood. Dip into a second a deflagrating spoon containing a small piece of burning phosphorus, and into a third dip a spoon with burning sulphur. The flames will be extinguished in all cases, showing the very inert nature of the gas. On withdrawing the spoon with the phosphorus it may reignite in the air. The reaction by which the nitrogen was secured in this case is somewhat complex, but by observation of certain details it may be understood. In the progress of the de- composition it will be noticed that vapor of water is given off along with the nitrogen gas, as it condenses in the de- FIG. 14. livery tube. At the end of the experiment, after cooling the tube, its contents may be shaken out and examined. In place of the red substance taken we have now a green pow- der, which is found to be partly insoluble in water. By mixing with water in a beaker, stirring a few minutes and filtering, something passes through the filter, leaving the green substance undissolved. By evaporating the filtered liquid we find a residue of potassium chloride, while the green substance the chemist recognizes easily as chromium oxide. We have, therefore, produced in the experiment, nitrogen gas, water, potassium chloride and chromium GENERAL CHEMISTRY. 109 oxide. Careful investigation shows that these substances are formed in the proportions illustrated by the following equation : K 8 Cr 2 O, + 2NH 4 Cl = N 2 + 2KC1 +4H 2 O +Cr 2 O 3 Potassium i Ammonium M -. _J_Potassium_l_ \i7_ t _ _l_Chromium dichroinate ' chloride ' chloride ' oxide. The combining weight of the potassium dichromate is 294.4, while that of the double part of ammonium chlo- ride taken is 107. We obtain 28 parts of nitrogen, or less than one-tenth the weight of the dichromate used. The process therefore is somewhat expensive. Many compounds containing nitrogen may be decom- posed, to liberate this substance, under proper conditions. In case of the compound taken, the ammonium chloride, it is necessary to add something to hold or fix the elements with which the nitrogen is here combined, and the di- chromate of potassium answers this purpose by furnishing potassium and oxygen to unite with these elements. Nitrogen has certain uses in the arts at the present time, but large quantities of the gas in the pure state are not required by them. General Properties. Nitrogen gas may be con- densed to a liquid by application of cold and pressure. The gas is but slightly soluble in water, and is character- ized by its extreme inertness or lack of positive properties. It is neither combustible nor a supporter of combustion, but it does unite directly with hot magnesium. A liter of the pure gas weighs 1.257 Gm. under standard conditions. THE ATMOSPHERE. The atmosphere is a mixture of oxygen and nitrogen, essentially, with smaller quantities of moisture, argon and carbon dioxide, and traces of other gases. The relation between oxygen and nitrogen has been suggested above, and can be shown by exact experiments. The amount of oxygen may be readily determined by the following, which 110 GENERAL CHEMISTRY. is merely a modification of our first experiment on the separation of nitrogen. Ex. 88. Procure a glass tube, sealed at one end, having a length of about 75 Cm. and an internal diameter of 1.5 to 2 Cm. If it is gradu- ated it will be so much the better for our purpose, but if it is not an approximate graduation may be made as follows : Pour a small meas- ured volume of water into the tube, when held in a vertical position and mark its level by means of a close fitting rubber ring shoved down the tube. Then add the same volume of water and mark the new level as before, and repeat the operation until the whole tube has been divided into equal small volumes. The rings should not be displaced by ordi- nary handling. Now pour about 20 Cc. of water into the tube and invert it in a deep jar of water. By means of a clamp on a lamp stand, fasten the tube in such a position that the levels of the water inside and outside are the same after the tube has stood long enough to have the air temperature. The tube should have a perfectly vertical position, Note the volume of the air enclosed with reference to the rubber rings. Next scrape off a piece of phosphorus, weighing two or three grams, and fasten it to a thin iron wire, a meter and a half in length, bent in the middle so that the two halves are ctose together. Now depress the wire with the phosphorus in the jar so that the limb with the phosphorus is brought beneath the opening of the tube. Then puil up the free end of the wireand guide the phosphorus with the fingers of the other hand so that it enters the tube. By means of the outside wire it can now be pulled up to the top of the tube, nearly, and there it should be left about 24 hours. In these manipulations, care must be taken not to bring the end of the tube above the surface of the water in the jar, and so change the air volume once read off. Great care must be taken to prevent the ignition of the phosphorus while handling it. This can be avoided by keeping it wet. When once up in the tube there is no further danger. The phosphorus undergoes slow oxidation and gradually combines with the oxygen present. By the end of 24 hours the reaction will be com- plete. Then shove down the wire and remove the phosphorus carefully and put it back under water. Observe that the water level in the tube is higher than before. By means of the clamp, and without touching the tube with the hands, depress it until the two water levels are the same again. Read off the gas volume now enclosed. The decrease in volume represents the oxygen only. Instead of inserting the phosphorus by the method described, the student may find it more convenient to close the end of the long tube, after noting the gas volume, by means of his finger, and then carefully lift it out of the water in the jar and dip it beneath the surface of water in a large bowl. The air volume remains unchanged. The tube can now be inclined to one side and the phosphorus, scraped and fastened to the wire in the same bowl, can be readily shoved up into the tube. If the wire is soft enough there will be no difficulty in bending it around the end of the tube so that the latter may again be closed by the finger and brought back into the jar and clamped. Before beginning the experi- ment, a piece of phosphorus of the proper diameter should be selected. GENERAL CHEMISTRY. Ill To obtain an accurate result by the above experiment it is essential that the temperature of the gas in the tube at the two readings remain the same, and also that the pressure of the air outside, or the barometric pressure, remain unchanged. These conditions are practically never attained and it is therefore necessary to make certain cor- rections to compensate for these changes. By increase of temperature gas volumes expand, and therefore, if the laboratory is warmer at the time of the second reading than at the first, the volume read off will be high and the loss (or the amount of oxygen) will be made to appear too low. If the temperature at ,the second reading is lower than before, the residual volume will be low and the oxygen will thus be obtained too high. This effect of temperature can readily be observed by the student by grasping the tube, still clamped in position, in the hand. The heat of the body, thus communicated to it, is suffi- cient to make a marked depression* of the water level in the inside of the tube. Before making readings the tube should, therefore, be handled as little as possible. Changes in air pressure outside change the inner volume also. An increase in the air pressure, indicated by elevation of the barometer, is communicated through the water and decreases the gas volume in the tube. The increased pressure forces the water down in the jar and therefore, because they are in communication, up into the tube. Following a decrease in barometric pressure the gas volume in the tube will expand. As preliminary to an explanation of the calculations of these corrections let the following experiment be made : Ex. 89. Suspend a thermometer in such a manner that its bulb hangs wuhin a few centimeters of the middle part of a graduated tube containing some air as in the last experiment. When the temperature appears "to be constant, and the volume of the gas therefore stationary, raise or depress the tube by means of the clamp until the water levels inside and outside are accurately the same. Now, read off the volume of the air as shown by the graduation, read the thermometer and the height of the barometer which should hang in the immediate vicinity of the other apparatus. The temperature within the tube is assumed to be the same as shown outside. With the water levels the same the pressure on the gas within the tube must be the same as that of the air, as 112 GENERAL CHEMISTRY. measured by the barometer. With these data at command let the student calculate the reduced volume at the assumed normal conditions of C. and an air pressure of 760 Mm., by the method explained below. REDUCTION OF GAS VOLUMES. Correction for Temperature. In comparing gas volumes it is necessary to refer them to* some standard temperature, which by common consent is C. As intimated above, all gases expand by increase in tem- perature, and practically at the same rate. By rate of ex- pansion or coefficient of expansion we understand the fraction of its volume at which it increases for an increase of 1 in temperature. This rate is the same for an increase from to 1 as it is for an increase from 10 to 11 or from 19 to 20 practically. In other words, it is constant, or so nearly constant that we assume it for our purpose. Let us represent the volume of a gas measured at by V and the volume of the same gas expanded to the tem- perature t, by V t . What is the relation of V t to V ? It is evident that this equation must be true, V t =V -[-increase. The increase in volume is made up of three factors, or is the product of three factors. One of these is the rate of expansion defined above, which we will call r. The second factor is the number of degrees of temperature through which the expansion takes place, and this we call /, while the third factor is the amount of V itself, or in other words is the number of units of volume in V . The actual increase in volume for two liters would evidently be twice as great as for one liter. Our equation therefore becomes, V (t , =V (0) +V (0) XrX/, or, by a slight alteration V (t) =V (0) (l+r/). GENERAL CHEMISTRY. 113 Now, as intimated, the rate of expansion of all gases is nearly the same and amounts to ^r of their volume at 0, for each degree of increase. Therefore, r=-yfa, or, ex- pressed decimally 0.00366. Making this substitution we have, V (t) =V (0) (l + 0.00366/). This is a fundamental equation and by transformation we get the next one, v V in t - (0) 1+0.00366/ In illustration of these equations assume that we have 100 Cc. of air at and wish to know its volume when warmed up to 25. That is, we wish to find V (tJ or V 20 . V (25) = 100 (1 + 0.00366X25) = 100 (1+0.0915) = 109. 15. The new volume is therefore 109.15 Cc. Conversely, if we have given this volume at 25 and wish to know what it becomes when cooled to we use the formula, V (0) __109.15 1.0915 = 100 Cc. As gases contract below zero at the same rate at which they expand above zero these formulas can be used for minus temperatures by change of sign. Correction for Pressure. We learn by experiment that the volume of a gas varies as the pressure changes, but inversely. That is, if the pressure becomes doubled the gas volume is contracted to one half. If the pressure decreases to one half the gas expands to fill double the 114 GENERAL CHEMISTRY. volume. If V represents the volume of a certain gas at the pressure, P, and V' the volume of the same gas at the pressure, P', it would follow from the above that VP = V'P' = V"P", etc. That is, the product of the volume and the correspond- ing pressure is a constant. We assume a standard pres- sure as we do a standard temperature in the measurement of gas volumes and this is usually taken as the average pressure of the air at the sea level, a pressure equivalent to that of a column of mercury 760 Mm. in height. If we let P = 760 Mm. and V the normal volume then it would follow that V _V'P' 760* That is, the reduced or normal volume is equal to the product of the observed volume and the observed pressure divided by 760 Mm., the standard pressure. In illustra- tion, if we read off a volume of 150 Cc. at 740 Mm. pres- sure, as shown by the barometer or other pressure gauge, the reduced volume must be = 146.05Cc. 760 The effect of changes of temperature and pressure are independent of each other; we can make, therefore, either correction first, and on this result calculate the other correction. Both corrections can be introduced in one formula. Let v represent the observed volume, h the observed barometric height and V the volume at and 760 Mm. Then y- v h (1 + 0.00366 /) 760 This formula is used where the gas is measured under GENERAL CHEMISTRY. 115 a pressure exactly equivalent to that of the barometric height, h. In practice certain modifications may be neces- sary. The calculation of the results of the last experiment is a case in point. The volume of the gas is read off over water; it is therefore saturated with moisture, which also exerts some pressure, even at low temperature. The air pressure, h> is therefore balanced by the true gas pressure, which we can call /, pJus the pressure of the aqueous vapor which we can call w. hp-{-w, therefore p=h w. Our formula above then becomes v (h #/) V = (1 -f- 0.00366 /) 760 The tension of water vapor is always expressed in mil- limeters of mercury, and can be found from a table in an earlier chapter. For three common temperatures it may be given here. t w 15... 12.7 Mm. 20 17.4 Mm. 25 , 23.5 Mm. Suppose now, that in the last experiment we read off a volume of 95.5 Cc. of air saturated with moisture, at a tem- perature of 20 C, while the barometer stood at 745 Mm. The reduced dry volume would then be found by substitu- tion as follows : V- 95.5 (74517.4) _69486 (1+0.00366X20) 760~815.6 In practice it may not be always convenient or desirable to depress the measuring tube until the levels inside and outside are the same. This is generally the case when 116 GENERAL CHEMISTRY. gases are measured and operated upon over mercury. A. column of liquid stands up in the tube and the pressure of this, also, must be brought into the calculation. The pressure of this column, like that of the aqueous vapor, must be subtracted from the height of the barometer as read off. If this column is mercury, measure its height in millimeters above the level of the mercury in the reservoir below and call it m. The general formula of reduction then becomes v (h m w). V = (1+0.00366 /) 760 In measuring gases over mercury it is generally best to add a minute drop of water to insure that they are fully saturated with moisture. Ex. 90. As an exercise in these calculations let the student meas- ure a gas volume over mercury in an accurately graduated tube, and make all reductions necessary by the last formula. A good barometer and a thermometer should be mounted near the gas tube. The exercise should be repeated until the principle involved is perfectly under- stood. Other Air Tests. The amount of oxygen in the atmosphere can be easily and quickly found, by absorbing it from a measured air volume, by means of an alkaline solution of pyrogallol. When great accuracy is desired rather elaborate apparatus must be used, but the method may be illustrated by a very simple experiment. Ex. 91. Select a long, narrow test-tube, which may be closed by the thumb, and by means of rubber rings, or otherwise, divide it into six equal divisions. Then, holding the tube vertically, pour in enough 10 per cent solution of potassium hydroxide to fill one division. Next incline the tube, and by means of a knife blade introduce about half a gram of dry pyrogallol in such a manner that it will rest on the side of the tube, but above the alkali. Then close the tube firmly by the thumb, and shake it thoroughly. After a minute, invert the tube, with the mouth under water, and remove the thumb. It will be seen that water enters to take the place of the oxygen absorbed by the pyrogallol and alkali. If the tube is allowed to cool to its original temperature, four divisions should remain filled with gas. The most accurate determination of the amount of GENERAL CHEMISTRY. 117 oxygen in the atmosphere is made by aid of the eudiometer as described in the chapter on water. A volume of air is measured in the eudiometer over mercury, the proper cor- rections being made. Hydrogen is then introduced, more than enough to combine with the oxygen present, and the new volume is measured. A spark is then passed through the gaseous mixture as described, and this causes the oxygen to unite with the requisite amount of hydrogen to form water. A contraction follows and after a time the new volume is read off and reduced to normal temperature and pressure. As two volumes of hydrogen unite with one of oxygen, one-third of the loss noted is the amount of the latter originally present in the reduced air volume. In 100 volumes of dry air there are 21 volumes of oxygen. OTHER CONSTITUENTS OF THE ATMOSPHERE. The amount of nitrogen in dry air was assumed, until quite recently, to amount to 79 volumes in 100, the traces of other gases being very minute. But in 1895 it was found by Rayleigh and Ramsay that a gaseous element resembling nitrogen is also present, and this element has always been included in what has been measured as nitrogen. Argon. This is the name which was given to the new element, but up to the present time its properties have not been fully described. The amount of argon in the air appears to be about 1 per cent. It is especially characterized by great inertness in power of combination, and this accounts for the fact that it has so long escaped recognition. Argon has been found elsewhere as well as in the atmosphere, and more recently a second gaseo-us element, termed helium, has been found, which frequently accompanies argon. The amount of this in the atmosphere is extremely small, but it is more abundant in the gases escaping from certain springs." Argon is left as a residue when a large body of air is 118 GENERAL CHEMISTRY. passed over hot copper for the removal of the oxygen and repeatedly over hot magnesium for the removal of the nitrogen. The nitrogen may also be removed by combin- ing it with oxygen by aid of the electric spark and over alkali to absorb the acid products formed. The excess of oxygen is afterward removed by the copper method or by pyrogallol, as described above. Carbon Dioxide. The air contains about 3 volumes of this gas in 10,000 under normal conditions. But in the atmosphere of crowded cities and in buildings it may be much increased. In the streets it is sometimes present to the extent of 6 or 7 volumes in 10,000, while in crowded rooms, poorly ventilated, it sometimes reaches 15 volumes in 10,000, or even higher. The gas is produced by proc- esses of respiration and combustion, and is absorbed from the air by the growth of vegetation. The total amount of this gas present in the air is enormously great, being estimated to exceed 3,000 billions of kilograms. Moisture, The amount of aqueous vapor present in the air varies within wide* limits, being largely dependent on the temperature. A cubic meter of air, if fully saturated at 0, contains 4.87 Gm. of moisture; at 10, under the same conditions, it contains 9.36 Gm.; at 15, 12.75 Gm.; at 20, 17.16 Gm.; at 25, 22.84 Gm., and at 30, 30.1 Gm. The air is seldom fully saturated, but often holds 75 per cent of this amount. Such an atmosphere is un- pleasantly moist. If the amount of moisture is below 50 per cent of that required for saturation the air appears dry to the skin. When an atmosphere nearly saturated with moisture is suddenly cooled to a temperature below that for which the water present is sufficient for saturation a part of this water must precipitate in the form of rain. On the other hand, if the air remains warm it may. hold a very large amount of moisture without precipitating, and through much of our summer weather in the United States this is often the case. Evaporation from the skin cannot take place if the air is already saturated with moisture, and such an atmosphere we describe as a "close" one. GENERAL CHEMISTRY. 119 The amounts of moisture and carbon dioxide in the air are best determined by aspirating a given volume of the air through a series of weighed absorbing tubes. The first of these tubes are charged with substances to take up the water; the following tubes contain something to absorb the carbon dioxide, usually caustic potassa. The increase of weight in the tubes, after passing the measured air volume, gives the amounts of the two substances absorbed. Ammonia. Traces of ammonia, NH 3 ,are found in the atmosphere at all times, and usually combined as am- monium carbonate. The ammonia is mostly derived from the decomposition of nitrogenous organic matter and although very small in relative amount in the whole vol- ume of air it is sufficiently great to be quite important. As it is carried down by the rain it enters the soil and there serves as a valuable food for growing plants. A large part of the nitrogen taken up by certain crops probably comes from the ammonia reaching the rootlets in this manner. Ozone. Traces of ozone, peroxide of hydrogen, and oxides of nitrogen are also often found in the air. These have been referred to already. Traces of sulphurous oxide and other gases are usually found in the air of cities and besides these, and of great importance, should be mentioned the small amount of organic dust everywhere present. This dust consists partly of living andpartly of dead matter. In the living mat- ter are included numerous minute microorganisms which are active in promoting fermentations and putrefaction. The dead organic matter comes largely from the decay and disintegration of animal and vegetable substances. CHAPTER VI. COMPOUNDS OF NITROGEN. /^COMBINATIONS of nitrogen formed by direct union ^* with the gas are rare, but a large number of sub- stances containing nitrogen can be made indirectly. Some of these are of great value and importance. NITROGEN AND OXYGEN. Nitrogen combines with oxygen to form five com- pounds, which are named as follows: Nitrogen monoxide N 2 O gas. Nitrogen dioxide NO or N 2 O 2 gas. Nitrogen trioxide N 2 O 3 volatile liquid. Nitrogen tetroxide NO 2 or N 2 O 4 volatile liquid. Nitrogen pentoxide N 2 O 6 solid. The production of the first and second of these will be, shown by experiment. Nitrogen Monoxide. This substance is known as laughing gas, or sometimes as nitrous oxide, and is readily made by the decomposition of a common crystalline sub- stance known as ammonium nitrate, as shown below. It is made on the large scale at the present time and is sold compressed in cylinders. Ex. 92. Dry a small flask holding about 250 Cc., and fit it with a perforated cork and long delivery tube bent so as to lead down from the flask, when mounted on a sand-bath, to a trough of water. Pour into the flask 10 to 15 Gm. of dry ammonium nitrate, insert the cork holding the tube, and support the flask on a sand-bath. Now apply heat, very GENERAL CHEMISTRY. 121 gently at first, which will soon melt the solid ammonium nitrate. Later, gas bubbles will escape from it, and it will appear to boil. Bring the lower end of the delivery tube beneath the surface of water in the pneumatic trough or basin, and collect several bottles of the escaping gas in the usual manner. As the bottles fill, remove them by aid of glass plates, and stand them on the table in the upright position. Test the gas by burning in it a splinter of wood, some charcoal, sulphur and phosphorus. For the last two use a deflagrating spoon. These sub- stances will burn almost as well as in oxygen. In performing this experiment observe certain precautions. The delivery tube should be wide, the heat should not be allowed to become higher than necessary to decompose the substance, and before removing the lamp the delivery tube should be withdrawn from the water for reasons already explained. The decomposition of the ammonium nitrate takes place according to this equation : NH 4 NO 3 =N 2 O + 2H 2 O Ammonium Nitrousi \yater nitrate ~ oxide 80 = 44 -f 36. 100 Gm. of ammonium nitrate yield, therefore, 55 Gm. of the oxide. One liter of the gas, at standard tem- perature and pressure, weighs 1.98 Gm., from which it follows that 100 Gm. of the nitrate will yield nearly 28 liters of the gas. Properties. At a low temperature the gas may be readily compressed to a liquid. It is slightly soluble in water, one volume of which dissolves about 1.3 volumes of the gas at 0. It forms no chemical combination in dissolving. At a moderately high temperature the gas is decomposed into its constituents, and this accounts for the fact that combustions follow so readily in it. If metallic sodium is strongly heated with a measured volume of the gas over mercury, combustion follows, and after the residue of nitrogen cools, it will be found to possess the volume of the original gas, from which it follows that the monoxide contains its own volume of nitrogen gas. Uses. As laughing gas it has been used for many years by dentists and surgeons for the production of mild anaesthesia. When employed for this purpose it must be 122 GENERAL CHEMISTRY. carefully washed by bubbling through water after leaving the generator. Nitrogen Dioxide is a substance formed in many reactions in the laboratory, and usually where nitric acid is decomposed by one of the heavy metals. Hydrochloric and sulphuric acids usually yield hydrogen, in contact with metals, but with nitric acid of about 1.2 sp. gr., nitrogen dioxide is liberated. This can take place only through complete decomposition of a part of the acid as shown in the next experiment. Ex. 93. Arrange a bottle as for generating hydrogen by action of dilute sulphuric acid on zinc, and charge it with about 20 Gm. of copper turnings and 25 Cc. of water. Pour in now through the funnel tube 25 to 50 Cc. of strong nitric acid. At first red fumes fill the bottle ; when these have been driven out connect with a receiving bottle over water, and collect as directed for hydrogen. This experiment should be tried where there is a good circulation of air, as the gas, or rather the product which it forms with oxygen, is very irritating. Collect several bottles of the gas, then fill the generator with water to check the reaction, throw away the acid thus diluted, wash and save any copper left. Test the gas as follows. Plunge a burning splinter in one bottle, and some burning sulphur in another. It will be seen that it does not support combustion. Remove the cover from a third bottle of the gas and bring the mouth of a clean, dry bottle over it. Holding the bottles to- gether invert them so that their contents will mix. Observe that red fumes are formed as the gas comes in contact with the air. The chemical reaction in this experiment is somewhat complicated, but has been determined by full investigations. It appears from these investigations, which cannot be given in detail here, that the decomposition of the acid by the metal takes place in several stages. But the final results are probably represented by this equation : 3Cu+8HNO 3 = 3Cu(NO 3 ) 2 +2NO+4H 2 O. A few pages in advance, when the properties of nitric acid are consfdered, some illustrations of the general behav- ior of this acid with metals will be given. The gas is often called nitrogen dioxide because of the fact that, for a given weight of nitrogen, it contains twice as much oxygen as the f rst oxide described. The older GENERAL CHEMISTRY. 123 view of its structure is represented by the formula, N 2 O 2 . It is now well known that this cannot be the formula. Properties. The gas is very slightly soluble in water and does not combine with alkali solutions. It may therefore be bubbled through a solution of sodium hydroxide to purify it in the method of preparation given above, which as described, does not yield a perfectly pure product. The gas combines with oxygen to form the substance NO 2 or N 2 O 4 , which was illustrated in the above experi- ment. The gas does not support the combustion of wood or sulphur, but if phosphorus, burning brightly, be plunged into it active combustion follows, because in this case the initial temperature is high enough to separate the oxygen from the nitrogen. Nitrogen Trioxide is not readily obtained in the pure state, but an illustration may be given of its formation. Ex. 94. Let the student mix about 10 Cc. of strong nitric acid with a gram of starch in a small flask. This is placed on a sand-bath in a fume closet and slowly heated. After a time red vapors appear in the flask and the reaction soon becomes violent. The lamp should then be removed. Dense red vapors escape from the flask. These consist of nitrogen trioxide with some tetroxide. In this experiment the starch decomposes the nitric acid, taking a part of its oxygen to form several complex bodies. The remainder of the nitrogen appears as triox- ide mainly. The trioxide dissolves in cold water, giving rise to a body known as nitrous acid : N 2 3 + H 2 =2HN0 3 Nitrogen _l_ w a tr Nitrous trioxide acid. Nitrous Acid in the pure state is not important as it is not stable. But it combines with alkalies, as caustic soda or caustic potassa, forming nitrites, which are very impor- tant substances for laboratory and manufacturing uses: = NaN0 2 +H 2 0. Nitrites are often made by the reduction of nitrates, 124 GENERAL CHEMISTRY. that is by the removal of oxygen from the latter salts. Such a reduction may be effected by fusion with lead, as illus- trated by this equation : All nitrites are readily soluble in water. Nitrogen Tetroxide. As explained above, this body is formed by the union of nitrogen dioxide with oxygen. The actual composition of the substance varies with the tem- perature, consisting mainly of N 2 O 4 at the ordinary work- ing temperature of 25C. At a few degrees below this point it may be condensed to a liquid. Nitrogen tetroxide is decomposed by contact with water, forming nitric acid and nitrogen dioxide. In presence of air, nitric acid is finally the sole product. In pure condition it is best made by the decomposition of lead nitrate by heat. The reaction is illustrated by this equation: By passing the gaseous products of the reaction through a TJ tube, immersed in a cooling mixture, the N 2 O 4 may be condensed to a yellow liquid and thus separated from the oxygen. This liquid boils at 22, about. Nitrogen Pentoxide is a laboratory product of no importance in the pure state. It is sometimes called nitric anhydride, and when dissolved in water forms nitric acid, as here illustrated. N 2 5 + H 2 0=2HN0 3 It is a white, crystalline product, and is usually made by extracting water from nitric acid by means of phosphoric anhydride, P 2 C>6- ^ IS not sta ble and breaks up into oxy- gen and nitrogen tetroxide. GENERAL CHEMISTRY. 125 NITRIC ACID. Occurrence. This acid is formed in traces in nature when electricity passes through moist air. In combined condition it is found as calcium, sodium, or potassium nitrate. Nitrates are produced in nature by a number of oxidation processes, which take place in the soil and in water and which are of the highest importance, as will be more fully explained some pages in advance. History. The acid seems to have been first made by distilling a mixture of saltpeter, alum and blue vitriol, and this as early as the ninth century. Later, in the middle of the seventeenth century, it was made by Glauber by a process similar to that still employed, that is, by distilling saltpeter with sulphuric acid. The actual composition of the acid was a subject of lively discussion until after the days of Lavoisier and Cavendish. The latter finally gave the true explanation of its formation and structure. Preparation. As mentioned, this substance is pro- duced in small traces by certain natural agencies, but in quantity it is always made by the decomposition of some nitrate. We have seen that common salt, or sodium chloride, is decomposed by sulphuric acid, yielding, as one product, hydrochloric acid. In a similar manner saltpeter, or the nitrate of potassium, is broken up by distilling it with sulphuric acid, yielding as the important product nitric acid. We can best illustrate this by the following experiment : Ex. 95. Arrange a glass retort with a flask as a receiver as shown by the next figure. Charge the retort with 25 to 30 Cc. of strong sul- phuric acid and about 50 Gm. of powdered potassium nitrate. Mount the retort on a sand-bath and apply heat, gently at first and then with the full gas pressure. The contents of the retort become thin and give off reddish vapors which pass over into the receiver, kept cold by water, and condense to a reddish yellow liquid; collect 10 to 15 Cc. of the acid liquid and apply tests to it as follows : Transfer it to test-tubes and in one dip a pine splinter. After a time note the appearance of the wood on its withdrawal. Add a little starch to the same tube and heat 126 GENERAL CHEMISTRY. gently (do this in a fume closet). Observe the decomposition already explained. In a second tube add some copper turnings to the acid. Observe the gas given off, and the color of the solution formed. To the acid in a third tube add some strong solution of ferrous sulphate, green vitriol, and observe the brown color formed, which disappears rapidly at first, but later remains as a ring in the middle of the liquid column. The liquid residue in the retort solidifies on cooling and then is removed with some difficulty. It is therefore f^AHKUfi <9-Qtl. FIG. 15. better to pour it through the tubulure of the retort, while still hot, into a porcelain dish. The retort is then allowed to cool and what remains can be readily washed out. The substance in the dish should be heated on a sand- bath, in a fume closet, as long as it continues to give off vapors. What is left soon becomes solid on cooling and is easily recognized by the chemist as potassium sulphate. We have therefore here a reaction in which potassium nitrate gives place to potassium sulphate, and free nitric GENERAL CHEMISTRY. 127 acid takes the place of the sulphuric. By the use of our symbols we can illustrate this change as follows: H 2 SO 4 -h2KNO 3 =K 2 SO 4 -h2HNO 3 Sulphuric_L_ Potassium -Potassium _j_ Nitric acid nitrate. sulphate ' acid. Nitric acid, it appears, is formed by the decomposition of a nitrate, and conversely, a nitrate can be made by the action of nitric acid on certain other substances, in illus- tration of which make the following experiment : Ex. 96. Dissolve about 5 Gm. of potassium carbonate (pearlash) in a little water and add to it some nitric acid, a little at a time. Effer- vescence begins as the liquids mix. Add the acid slowly as long as gas escapes after thoroughly stirring the solution. Then add a few drops more of the acid and evaporate the liquid resulting in a porcelain dish to complete dryness, stirring well at the end of the operation. Compare the solid substance left, in appearance and taste, with the nitrate used in the last experiment. From what has been given above it is evident that nitric acid is a very strong and corrosive substance. In the pure form it is not very stable, and therefore it appears in commerce diluted with water. It could not be safely handled under other conditions. On the large scale the decomposition of the saltpeter is carried out in iron retorts or boilers holding sometimes tons of the raw materials. Sodium nitrate or Chili salt- peter is now commonly used instead of potassium nitrate, as it is much cheaper. The acid fumes which distill over are collected in a series of well glazed earthenware Woulfe bottles, which contain a little water. It is not practicable to make or handle absolute nitric acid. The commercial product always contains some water. Several grades are sold having specific gravities usually from 1.38 to 1.45, corresponding to acids of 61 to 77 per cent strength ap- proximately. A strong red acid is found in the market, known as fuming nitric acid. This acid contains oxides of nitrogen dissolved. Nitric acid is most characteristically distinguished by its power of oxidation or supporting a kind of combustion. Indeed, there are cases in which it may be made to give 128 GENERAL CHEMISTRY. up oxygen directly to ordinary combustible substances, as illustrated by the following experiment : Ex. 97. Place a small beaker on an iron dish containing sand. Pour 20 Cc. of strong fuming nitric acid into the beaker. Grasp a rod of charcoal, about 5 Cm. long and 5 Mm. in thickness, with iron forceps at one end, and ignite the other end in a gas flame. When it is glowing brightly dip the burning end beneath the strong acid in the beaker and observe that the combustion continues with evolution of red fumes in quantity. The oxygen of the acid is in part taken by the charcoal, while the rest is held by the hydrogen and nitrogen. In performing this experiment care must be taken to avoid touching the glass with the hot charcoal. As this might break it and spill the acid, the sand is placed beneath the beaker to catch the acid if this accident happens. The experiment must be made in a fume cjoset, as the gases given off are very offensive. In any experiment with nitric acid in which red fumes are abundantly given off we may be certain that oxidation is taking place. That is, a part of the acid, at least, is undergoing decomposition by which oxygen is given to some other substance. In the above case the charcoal burns at the expense of the oxygen, so furnished, as well as it does by the aid of atmospheric oxygen. This property of oxidation is not confined to the acid alone, but is found in marked degree in many of the com- binations of nitric acid, or nitrates, as can be shown by the following experiment with potassium nitrate : Ex. 98. Dry and powder about 10 Gm. of potassium nitrate. Pour it into an ordinary dry test-tube, and heat this carefully in the flame of the Bunsen burner. The nitrate soon begins to melt, and finally all of it becomes liquid. Continue to heat it carefully in the flame, moving the tube to and fro, to evenly distribute the heat. Then drop in a small piece of charcoal and observe that it soon ignites and burns with a hiss- ing noise. When the first piece is consumed add a second, which burns in the same manner, the tube being still held in the flame. Then drop in a small fragment of sulphur, which likewise burns at the expense of the oxygen taken from the nitrate. In performing this experiment the tube should be heated over a sand-bath to catch the liquid saltpeter in case of breakage, which, however, need not happen if care is taken. The above experiment illustrates in a very marked manner the oxygen furnishing power of a nitrate. Niter, either potassium nitrate or sodium nitrate, is used with sulphur and charcoal in the production of ordinary gun- GENERAL CHEMISTRY. 129 powder, which contains the three substances in certain proportions. Nitric acid is employed in the arts mainly on account of its oxidizing properties, rather than because of its acidity. A few metals resist the action of strong nitric acid, most of them are dissolved forming nitrates, while a few are converted into oxides without solution. This is trueof tin and antimony as will be shown later by experiment. The behavior of nitric acid in dissolving metals is not perfectly clear, and several theories have been advanced to account for it. To discuss these would be out of place in an elementary book. But one view which has long been advocated may be illustrated by considering the reaction between copper and nitric acid. It is assumed that this takes place in two stages. We have first, apparently, what may be called the normal reaction between a metal and an acid, that is, a solution of the metal with liberation of hydrogen, as follows: Cu + 2HNO 3 = Cu(NO s \ + H 2 Coppr-(- Nit F* c = Copper J-Hydrogen. acid nitrate This hydrogen, however, instead of escaping in the free state as it does from sulphuric acid and zinc, seems to react on the excess of nitric acid present and decomposes it. Hydrogen in this condition is called nascent hydrogen, and because of its powerful attraction for oxygen is capable of acting as a reducing agent. At any rate, it is absorbed by the acid, and decomposition products of the latter escape. The reaction here is probably represented by this equation : 3H 2 + 2HN0 3 = 2NO +4H 2 O Hydrogen+ Nitric =^ff+ Water. By combining the two equations we have the one given some pages back : 3Cu+8HNO 3 =3Cu(NO 3 ) 2 +2NO+4H 2 O. According to other views the reaction is possibly ex- pressed by these equations: 130 GENERAL CHEMISTRY. 2HNO 3 + 3Cu = H 2 O+3CuO-f2NO CuO+2HNO 3 = Cu(NO 3 ) 2 + H 2 O. With certain metals and dilute nitric acid this assumed reduction by nascent hydrogen seems to go much further, leaving ammonia as the final product. Dilute nitric acid and zinc seem to react on each other in this way : 4Zn+8HNO 3 -4Zn(NO 3 ) 2 +4H 2 4H 2 + HNO 3 =NH 3 -f3H 2 O. By combination we have finally : 4Zn+9HNO s =NH s +4Zn(NO s ) g +3H 8 O. The ammonia remains combined with the excess of acid as ammonium nitrate. Physical Properties. Absolute nitric acid has the specific gravity of 1.53 at 15. It mixes with water in all proportions and decomposes on heating. The pure strong acid begins to boil at about 86 but decomposes into water, oxygen and nitrogen tetroxide. Under the normal pressure an acid of 68 per cent strength distills without change. The concentrated acid does not dissolve iron and may therefore be shipped in iron drums. Uses. Nitric acid is employed in many ways. A great deal is used in making nitro-celluloses and nitro-glycerol for explosives. Much is used also in making nitro-benzene and allied bodies employed in the color industries. Many nitrates are made by the solution of metals, carbonates or oxides. These nitrates are, in some cases, employed as oxidizing agents. With hydrochloric acid, nitric acid yields a valuable solution known as aqua regia, referred to below. Nitrates. The nitrates found in many places in nature are produced usually by a series of oxidation processes from animal or vegetable matter containing nitrogen. This matter reaches the soil often in the form of animal waste GENERAL CHEMISTRY. 131 or excreta. The changes which it there undergoes by which its nitrogen becomes combined with oxygen are largely the result of the action of minute microscopic vegetable cells known as bacteria. Many of these bacteria have the power of decomposing nitrogenous organic mat- ter in such a manner that in presence of oxygen the nitro- gen becomes united to it. Soils become enriched in this manner, as the nitrates are among the best foods for grow- ing plants. In many cases it is probable that the formation of ammoniacal compounds from more complex organic substances precedes the oxidation stage or production of nitrites and nitrates. This is usually the case in the de- composition of sewage in streams. Urea and other organic bodies are broken down through bacterial agency and ammonia results. In presence of sufficient air this ammo- nia later gives place to nitrites and nitrates. If this nitri- fication in the soil takes place in the presence of calcium bicarbonate, which is a common condition, calcium nitrate results. This salt is very soluble and may be carried through the soil to appear later on the sides of caverns as cave niter. The deposit in the Mammoth Cave is an illustration. Certain nitrates often appear as an efflorescence on the surface of soil in hot countries. In India such efflorescence, mainly potassium nitrate, is abundant enough to be collected as an article of commerce. The great beds of sodium nitrate in western South America were doubtless produced by the oxidation of decaying marine vegetation. These deposits furnish a large part of the niter in com- merce to-day. Aqua Regia. This is a mixture of 2 to 3 volumes of strong hydro- chloric acid with 1 volume of nitric acid. A decomposi- tion takes place by which two products, known as nitrosyl chloride, NOC1, and nitroxyl chloride, NO 2 C1, are formed. This mixture has strong solvent properties, as metals and ores may be dissolved in it which cannot be dissolved by either nitric or hydrochloric acid alone. It is therefore a valuable reagent in the laboratory, as will appear later. 132 GENERAL CHEMISTRY. NITROGEN AND HYDROGEN. AMMONIA. Although several compounds of hydrogen with nitrogen are known, only one of them, ammonia, is technically im- portant. This substance occurs in small amounts in air and water, but usually in combination with something else. History. Combinations of ammonia were known to the alchemists. One of the most important of these, known as sal ammoniac, was obtained by many different processes, and finally by distillation of animal refuse. The carbonate of ammonium resulted from this distillation, and this on treatment with hydrochloric acid furnished the chloride or sal ammoniac. In 1774 Priestley found that this substance when distilled with lime yields a gas, ex- tremely soluble in water and which may be easily decom- posed by the electric spark. The composition of this gas was determined by Davy and others. It is represented by the formula, NH 3 . Preparation of Ammonia. Ammonia in the free gaseous condition has been made by the direct combina- tion of its elements through the aid of the electric spark. But this preparation has no practical importance. On the large scale, and in laboratory experiments, it is best made by the decomposition of certain of its compounds, called salts of ammonium. This is illustrated by the next experi- ment. Ex. 99. Arrange a glass flask and Woulfe bottles as in the produc- tion of hydrochloric acid, some pages back. Each Woulfe bottle should contain about 100 Cc. of distilled water. In the flask mix about 30 Gm. of ammonium chloride and the same weight of slaked lime. Add enough water to make a thick liquid mixture on shaking. Close the flask, make the connections and apply a gentle heat, which may be gradually increased. A gas is given off from the heated mixture, which passes over and collects in the water of the first Woulfe bottle mainly. Some reaches the second bottle, and little or none the small flask. After the application of strong heat during half an hour, detach the bottles and test their contents as given below. Ex. 100. Take half a test-tube full of the liquid from each of the two Woulfe bottles and the small flask, and add to each 2 drops of an GENERAL CHEMISTRY. 133 aqueous solution of methyl orange (1:1000). This solution imparts a yellow color with alkalies, which is characteristic. In the liquid from the first Woulfe bottle the reaction should be strong, but much weaker in the others. Add dilute hydrochloric acid now to each test-tube, a drop at a time, and observe that in the first case many drops may be necessary to change the color from yellow to pink, but that in the sec- ond and third very much less is necessary. Acids impart a pink color to solutions of methyl orange. Next repeat the same experiment, using 10 drops of a weak alcoholic solution of phenol-phthalein(l:1000) instead of the methyl orange. A deep crimson red color is now obtained with the first test-tube, and weaker shades with the others. When hydro- chloric acid is added these colors disappear. Alkalies in general give a red color with phenol-phthalein, but in acids there is no color reaction. These experiments prove the strong alkalinity of the ammonia solution. *'1G. 16. Ex, 101. Thoroughly clean three small porcelain evaporating dishes. In one pour about 10 Cc. of the ammonia solution from the first Woulfe bottle. In another take about the same volume of dilute hydrochloric acid, while in the third equal volumes of the ammonia so- lution and dilute hydrochloric acid are mixed. During the mixing a great volume of white fumes is produced. Evaporate the three solutions slowly on a sand-bath. In the case of the first no residue will be left. The same is true of the second, or hydrochloric acid solution, showing the complete volatility of both of these products. But in the third case, with the mixture, we have left a white residue, which is identical with the ammonium chloride used in the experiment on the production of ammonia. It appears, therefore, that while ammonia and hydrochloric acid are extremely volatile, the product of their union is not, or at any 134 GENERAL CHEMISTRY. rate, but slightly at the temperature employed. That it is volatile at a higher temperature is shown next, by heating the dish containing the residue more strongly. Dense white fumes are given off, leaving practically nothing in the dish. All of these experiments must be made in the fume closet. The above experiments illustrate the method by which ammonia is obtained on the large scale. Crude ammonium chloride, produced as a by-product in the manufacture of illuminating gas, is distilled with slaked lime, and the gas given off is collected in water. An impure ammonia water is thus made, which is saturated with hydrochloric acid yielding ammonium chloride, not yet pure, but much better than the first crude substance. As this substance is vola- tile it may be greatly purified by sublimation. The sub- limed salt is sent into commerce and used for many pur- poses. If this sublimed salt is heated again with slaked lime a very nearly pure ammonia gas is given off which may be absorbed in distilled water, yielding a concentrated am- monia solution. By neutralizing this with pure hydro- chloric, nitric or sulphuric acid we obtain, on evaporation, pure chloride, nitrate or sulphate of ammonium. In many large chemical works at the present time pure ammonia water is obtained in one operation from gas liquor. Soft coal containsa little nitrogen, and in the man- ufacture of illuminating gas by the distillation of such coal the nitrogen becomes converted into ammonia which is carried along with the gas until a large washing tank, called the hydraulic main, is reached. Here the very sol- uble ammonia dissolves and combines with carbonic acid and hydrogen sulphide from the gas, to form carbon- ate and sulphide of ammonium. The water in this hydraulic main has to be frequently renewed. To recover the am- monia from it, it is run into large boilers, mixed with slaked lime and distilled. The lime decomposes the ammo- nium salts, setting free ammonia gas, which passes through a series of cooling pipes and small washing reser- voirs and is then absorbed in distilled water. GENERAL CHEMISTRY. 135 The changes referred to above may be represented by equations: Ca0 2 H 2 +2NH 4 Cl = CaCl 2 +2NH 3 +2H 2 O Slaked _l Ammonium Calcium_|_ A _|_ w , lime chloride chloride ^Ammonia -(-Water NH 3 + HC1 = NH 4 C1 Hdbl Ammonia+ ^tric ^Ammonium 2NH 3 +H 2 S0 4 = (NH 4 ) 2 S0 4 Ammonia+Sulphuric = Ammonium Finally, we sometimesi speak of the solution of am- monia in water as ammonium hydroxide, NH 4 OH, which is NH 3 -{-H 2 O. If we may assume that this body actually exists in solution, then our reactions should be written after this manner: NH 4 OH+ HC1 = NH 4 C1+H 2 O Ammonium _|_ Hydrochloric Ammonium_|_ w . hydroxide ~ acid " chloride r waier ' We have here a behavior analogous to that by which we obtained sodium chloride from hydrochloric acid and sodium hydroxide, or caustic soda. The solution of am- monia gas in water possesses the properties of sodium and potassium hydroxides in a marked degree, as will be shown later. The term caustic ammonia has been sometimes ap- plied to the solution. Other important properties of the gas remain to be shown. For this purpose we may decompose more am- monium chloride with slaked lime, but the gas may be obtained more conveniently by heating the strong solution, or commercial ammonia water. The following experiment will illustrate this: Ex. 102. Fit a flask holding about 300 Cc. with a perforated stop- 136 GENERAL CHEMISTRY. per, through which pass a straight glass tube about 20 Cm. in length. Pour about 50 Cc. of the strongest ammonia water of the laboratory into the flask, close it as explained, and then support it on a sand-bath in a stable position, with a small ring of a lamp stand around the neck of the flask, for instance. Then apply heat. This will cause the ammonia gas to escape from the solution and pass out through the tube. When the escape of gas becomes rapid as shown by the apparent boil- ing in the flask, hold a dry bottle over the open end of the tube, mouth downward, so that the gas may enter and force the air out. As the gas is but little more than half as heavy as air, it may be collected readily in this manner. The tube should reach nearly to the bottom of the in- verted bottle. After a few minutes the escape of the gas into the air will show that the bottle is full. Lift it up carefully and close the mouth with a glass plate. It m-ay be then placed, mouth still downward, on a table, while a second bottle of the gas is being collected in the same man- ner. Remove the second when full, and collect finally a third, giving to this more time than to the others. Remove this bottle, close it quickly with a glass plate, and bring the mouth of the bottle, still held downward, under water in a basin. Then remove the plate, when it will be noticed that the water rushes up and nearly or quite fills the bottle. Had the air been quite expelled in the collection of the gas, the water should completely fill the bottle, showing the quick absorp- tion of the gas by the water. Apply a flame test to one of the other bottles. To this end lift the bottle from the table, mouth still down, and insert a burning taper. It will be extinguished and no flame will re- main at the mouth, showing that the gas is neither combustible as is hy- drogen, nor a supporter of combustion, as oxygen. With the third bot- tle of the gas make this experiment: Lift it from the plate and push up into it a strip of perfectly dry red litmus paper. The change of the color to blue will not be rapid. Then moisten a piece of red litmus paper in fresh water which has not been exposed to the ammonia fumes, bring this to the bottle and observe that the change to blue is immedi- ate. The dry gas is not an alkali but becomes so in the presence of moisture. Properties. At the temperature of C. and under a pressure of 760 Mm. water absorbs about 1,150 times its volume of the gas. A cubic centimeter of water absorbs, therefore, over a liter of the gas. At 30 nearly 500 volumes are absorbed by one volume of water. It will be seen in what follows that ammonia solution is a reagent of great value in the laboratory, as it is employed for many pur- poses. The gas is also readily soluble in alcohol and the solution so made has several applications. At a low temperature dry ammonia gas may be readily condensed to the liquid form and is then usually called anhydrous ammonia. At a temperature of 15 a pressure of GENERAL CHEMISTRY. 137 about seven atmospheres is required for the condensation. This liquid boils at a temperature of 38. Like all con- densed gases, anhydrous ammonia absorbs a large amount of heat on passing from the liquid to the gaseous condition, and advantage is taken of this fact in refrigerating or in the production of ice. To accomplish this the condensed ammonia is allowed to expand from strong storage tanks into a system of pipes. These pipes may be arranged around the walls and ceilings of rooms to be cooled, or they may be built in more compact form and immersed in a brine reservoir. In this case the brine becomes cooled to a low temperature and if it is pumped into pipes it may be made to circulate through rooms or buildings where cooling is desired. The cold brine reservoir may also be used in the production of artificial ice. It is simply necessary to immerse tanks of distilled water in the brine and allow them to remain there a day or more. The water freezes to a block of clear, pure ice, which is easily removed. The brine for this purpose must have a temperature of 5 to 10 centigrade. The expanded ammonia is compressed again by powerful pumps and so used continuously. One liter of ammonia gas under standard conditions weighs 0.765 Gm. The specific gravity referred to air is 0.589; it is therefore one of the lightest gases known. It does not support combustion and can be burned only with difficulty. The composition of the gas may be easily deter- mined by decomposing it with the electric spark. This yields one volume of nitrogen to three volumes of hydro- gen, which may be shown in a eudiometer similar to the one used in the analysis of water. Hydroxylamin. This is a substance which may be looked upon as am- monia, NH 3 , in which one atom of hydrogen is replaced by the group, OH, called the hydroxyl group. It is an un- stable crystalline compound whichdecomposes very readily. Its solution in water is more stable and the salts, which may be compared with the ammonium salts, are easily made and preserved. The hydrochloride, NH 2 OH.HC1, 1&8 GENERAL CHEMISTRY. is the most important. A reaction by which it is often made depends on the reduction of the gas NO by nascent hydrogen. 2NO+3H 2 =2NH 2 OH. Hydrazin. This is a compound having the formula N 2 H 4 and made by decomposition of certain complex organic compounds. It is not stable in the free condition, but occurs as a sul- phate, N 2 H 4 H 2 SO 4 . This sulphate may be decomposed by alkalies yielding a hydrate, N 2 H 4 H 2 O, which is a liquid resembling ammonia in some properties. With acids this hydrate yields well defined salts. Some organic derivatives of hydrazin are bodies of great practical importance. Hydronitric Acid. This is a peculiar compound recently discovered and studied having the formula HN 3 . It is best made by this series of reactions : ammonia gas is led over sodium in a heated porcelain tube, yielding sodium amid, NaNH 2 : Na-fNH 3 =NaNH 2 + H. When this reaction is complete a current of dry nitrous oxide is passed through the tube, acting on the amid in this way : The sodium nitride, NaN 3 , distilled with dilute sul- phuric acid yields free hydronitric acid, or azoimid. In the pure condition the acid is a clear mobile liquid which explodes spontaneously with great violence. In water so- lution it is more stable. It forms salts with most of the metals, on some of which it has a marked solvent action, even attacking gold. It is a poison and destroys the skin rapidly. The odor of the acid is extremely disagreeable. GENERAL CHEMISTRY. 139 When inhaled the vapor produces violent headache. The salts are called nitrides and in many respects they resemble chlorides. The structure of the acid is probably N H-N NITROGEN AND THE HALOGENS. Nitrogen forms compounds with chlorine, bromine and iodine, which are all very singular in this respect that they are explosive to a high degree. Of these the iodine com- pound is most easily made and with safety. Ex. 103. In a small porcelain dish rub about half a gram of iodine to a fine powder and cover it with a few cubic centimeters of strong ammonia water. Stir repeatedly with a glass rod and after half an hour wash the contents of the dish into two filters. This residue consists of a dark powder, commonly called nitrogen iodide. Wash it on the filters with a little alcohol to remove any unchanged iodine, and then displace the alcohol by washing with water several times. Remove the filters from their funnels and hang them up to dry. It will be found that when the product is perfectly dry the slightest agitation is sufficient to explode it, with a sharp report. There has been some uncertainty regarding the com- position of these explosive bodies, and it appears that under different conditions different products are obtained. What is commonly called nitrogen iodide is probably a mixture of NI 3 and NHI 2 . The chlorine compound is a yellowish liquid having the composition NC1 3 , probably. FURTHER THEORETICAL CONSIDERATIONS. In the third chapter an outline of the atomic theory was presented and it will now be in place to introduce other points of a theoretical nature, as the student has become familiar with the preparation and properties of several new and important substances. The conception of atoms and molecules has been explained, and it was shown that the combination of atoms 140 GENERAL CHEMISTRY. in these molecules takes place in definite and fixed propor- tions. In all cases 23 parts of sodium combine with 35.5 parts of chlorine ; 39.1 parts of potassium combine with 35.5 parts of chlorine. Besides these no other combina- tions of sodium, potassium and chlorine are known. In these cases an atom of one element combines with an atom of the other. An atom of hydrogen weighing 1 combines with an atom of chlorine weighing 35.5 and the result is a molecule of hydrochloric acid with a weight of 36 5 on our arbitrary scale. But when we come to consider the com- bination of oxygen with hydrogen we find that the amount of the latter which unites with the atomic weight or 16 parts of oxygen, to form water, is just twice as great as the amount which combines with 35.5 parts of chlorine. The oxygen atom appears to have, therefore, double the com- bining power of the chlorine atom. The amount of hydro- gen which combines with one atom of nitrogen to form ammonia is three times that which combines with 35.5 parts or one atom of chlorine, and the amount of hydrogen which combines with one atom of carbon to form a com- pound known as marsh gas, or methane, is just four times that which will combine with one atom of chlorine. The combining powers or valencies of the ultimate atoms are therefore different and a study of the whole number of atoms known shows that some resemble hydrogen or chlo- rine, and these are called univalent; some resemble oxygen and are called bivalent; some resemble nitrogen and are called trivalent; some resemble carbon and are called quad- rivalent, while other atoms have still higher powers of combination. Of the real nature of this valence we know nothing and our methods of representing and describing it are at best crude. For sake of simplicity in writing formu- las where the valencies of the atoms are expressed we make use of dashes as in the following figures: H / ! i H , 0, N, C , H 0-H, H C H 1 GENERAL CHEMISTRY. 141 But there are many cases in which elements combine with each other in more than one proportion, cases in which, apparently, there are multiple combining weights. It will be shown that two oxides of carbon may be pre- pared. One of these is known as carbon monoxide, and the other as carbon dioxide. By exact analysis it has been shown that the ratio of the carbon to the oxygen in the first of these is 1 : 1.333, while in the second it is 1 : 2.666. That is, we have for a given weight of carbon twice as much oxygen in one case as in the other. Two oxides of nitrogen were prepared, and it was explained that three others are known. An interesting fact was shown by the analyses which chemists made of these. The ratios of nitrogen to oxygen in the five compounds are given in the following table : N O 1st 1 0.5714 2d 1 1.1428 3d.. ..!.. ..1.7142 5th 1 2.8571 The proportions of oxygen in these compounds stand to each other in the relation, 1, 2, 3, 4 and 5. Two oxides of sulphur are known also, in which the amounts of sulphur present stand to each other in the exact relation of 2 to 3. In investigating fully these three classes of compounds it soon becomes apparent that the union of oxygen with nitrogen, carbon or sulphur takes place always as here rep- resented, and such facts give the strongest support to the atomic theory. But analysis shows and the above table illustrates an- other point. It appears that nitrogen, under different con- ditions, has different capacities for holding or combining with oxygen, and we express this by saying that the valency of the nitrogen varies. Nitrogen is not the only element in which the valency or capacity for combination is variable. This seems to be, indeed, the rule rather than the exception and an attempt is made in the accompanying table to show the variations in valency in the more important elements : 142 GENERAL CHEMISTRY. TABLE OF VALENCY. Name. Symbol Valence. Aluminum Al Sb As Ba Bi B Br Cd Ca C Cl Cr Co Cu F Au H I Ir Fe Pb Li Mg Mn Hg Mo Ni N O P Pt K Se Si Ag Na Sr S Te Tl Sn Ti W U Zn Zr II I II II II I II II II I I I I II II 11 I II II II II II I II II I II I I II II 11 I II II II III III III III III III III III III III III III III III III V V V V IV V IV VI IV V IV IV IV IV IV VI IV V V IV IV VI IV IV VI IV VI IV IV IV VI IV VI IV VII VII VII (VII) Antimony Arsenic Barium Bismuth Boron Bromine Cadmium . Calcium Carbon Chlorine Chromium Cobalt CoDDer Fluorine . . Gold Hydrogen Iodine Iridium Iron Lead Lithium Magnesium Manganese Mercury Molybdenum Nickel Nitrogen Oxygen Phosphorus Platinum . Potassium . . . Selenium . Silicon Silver Sodium Strontium Sulphur . Tellurium ... Thallium Tin Titanium Tungsten . . Zinc Zirconium GENERAL CHEMISTRY. 143 Under the head of the compounds of chlorine a list of acids formed by the union of chlorine with hydrogen and oxygen was given. The names and formulas of these acids are here repeated and by means of dashes the variation in the valency of the chlorine is indicated. H O Cl hypochlorous acid. H O Cl=zO chlorous acid. H O Cl chloric acid. H O C1=O perchloric acid. In the above formulas the hydrogen has always a val- ence of one and the oxygen a valence of two. The formu- las indicate that the oxygen by its double combining power links the hydrogen to the chlorine. In the first formula the chlorine has a valence of one; in the second of three; in the third of five; in the fourth of seven. Through this increase in valence the chlorine is able to hold more and more oxygen, but, as already suggested, of the nature of this valence and the reasons for its variations we know nothing. Graphic Formulas. Formulas written like the above with the symbols detached and separated by dashes are called graphic or structural formulas. It is intended to represent by them the manner of combination of the atoms, that is to show which are linked or joined in the molecule. Formulas in which no attempt is made to show structure or the mode of combination are called empirical formulas. For economy of space these are commonly employed, but graphic formulas are exceedingly valuable and in many case's almost indispensable in clearly representing reac- tions. Their widest application is found in explaining the structure of complex organic compounds, but even in rep- resenting comparatively simple substances their use is 144 GENERAL CHEMISTRY. apparent. Sulphuric acid is commonly represented by the formula H 2 SO 4 , but experiment shows that the atoms of hydrogen are linked to the sulphur by aid of oxygen. We indicate this view then, by the more complete formula: H O O=0 The nucleus atom of sulphur is shown here as having a valence of six. Each oxygen atom has a valence of two and each hydrogen atom a valence of one. Such formulas will be frequently used in what follows but the student must remember that all such attempts to indicate structure are at best insufficient. Molecules are formed by aggregations of atoms in space of three dimensions, and not in two dimensions as we are for convenience obliged to show them, and in no case, as yet, is our knowledge accurate enough to indicate satisfac- torily the space relations of these atoms, beyond what is suggested above for sulphuric acid. The acids referred to above are qualitatively alike; they all contain hydrogen, chlorine and oxygen. To distinguish between them we employ certain prefixes and terminations, and the student will observe that the same are used in many similar cases. These are: hypo ous. ous. ic. per ic. It will be remembered that we have nitrous and nitric acids, and later hyposulphurous, sulphurous and sulphuric acids, hypophosphorous, phosphorous and phosphoric acids, and others distinguished in the same manner will be mentioned. The four chlorine acids, the sulphur acids, and the phosphorus acids differ exactly in the same man- ner, and that is in the valencies of the characteristic ele- ments present, and consequently in the amount of oxygen held. The acids having the lowest amount of oxygen are in GENERAL CHEMISTRY. 145 all cases designated as hypo ous or ous acids, while those with more oxygen or with greater valency are called ic acids. The prefixes hypo and per (or hyper) are from Greek prepositions meaning respectively less than, under, below, and more than, over, above. The terminations ous and ic are arbitrary indications of less or greater valence. A hypo acid is therefore one in which the characteristic element (the chlorine, sulphur, phosphorus, etc.) has a lower valence than it has in the ous acid. In the ous acid the valence is less than in the ic acid, and in the ic acid it is lower than in the per ic acid. It has been stated already that a salt is formed by re- placing the hydrogen of an acid by a metallic atom. Salts formed from acids containing but two elements, binary acids, take the termination ide. Thus, from hydrochloric acid, HC1, we obtain chlorides, NaCl, KC1, FeCl 2 and others. From hydrobromic acid we have bromides, from hydriodic acid we have iodides. On the other hand, if we consider the so-called ternary acids, those with three ele- ments, as were illustrated above, we have, on replacement of the hydrogen, salts corresponding to the acids. The nomenclature of the salts may be briefly indicated as follows : hypo ous acids yield hypo ite salts ous acids yield ite salts ic acids yield ate salts per ic acids yield per ate salts Thus, we have potassium chlorate corresponding to chloric acid, and sodium hypochlorite corresponding to hypochlorous acid. These designations are arbitrary but the student should make himself familiar with them as early as possible, as they are of frequent occurrence throughout the book. CHAPTER VII. SULPHUR AND IT5 COMPOUNDS, SELENIUM AND TELLURIUM. OULPHUR is an element which occurs abundantly in O nature in the free state and in many sulphides and sulphates. It is widely distributed. Preparation. Crude native sulphur occurs in Sicily in large quantities and is refined by very simple processes. Sulphur melts at a relatively low temperature and the first refining consists in melting it away from the accompany- ing earthy materials. The crude ore is heaped up in a large pit dug in a hillside and ignited. The combustion of a part of the sulphur furnishes heat enough to melt the rest which collects at the bottom of the pit and then escapes through an opening leading to a lower reservoir. The product so obtained is not pure but may be refined by distillation from large retorts. Some sulphur is obtained by distillation of certain sulphide ores containing it, but the amount so produced is not important. At a high heat a sulphide of iron decom- poses as here represented: 3FeS 2 =Fe 3 S 4 +S 3 . That is, one-third of the total sulphur maybe obtained in the free state. It will be recalled that the dioxide of manganese may be broken up in the same manner: = Mn 3 O 4 -fO 3 . Vast quantities of sulphur exist in deposits occurring GENERAL CHEMISTRY. 14? in southern Louisiana, at a considerable depth below the surface of the earth. After many futile attempts to mine this, it has recently been found possible to bring it to the top of the ground in this manner: Holes are bored down to the deposit, and these are doubly piped. A large pipe fills the boring, and inside of this a smaller one goes down to a somewhat greater depth. Hot water under great pressure and at a high temperature, about 170 C, is pumped down in the space between the pipes. This melts the sulphur in the deposit and ultimately forces it up through the inner pipe. To aid in maintaining a high temperature, a third small pipe passes down through the one which conveys the sulphur to the surface. A current of hot air is forced down in this, but the pressure on it is less than that on the water. The molten sulphur on reaching the surface is run into shallow pans to cool and solidify. It is nearly pure, and for most purposes needs no refining. Properties of Sulphur. Sulphur appears in com- merce in fine powder or "flowers of sulphur," and as roll sulphur or "brimstone." Both varieties are insoluble in water, but soluble in carbon disulphide, and in several other liquids. The solubility in carbon disulphide may be shown by experiment. Ex. 104. Take about 10 Cc. of carbon disulphide in a test-tube and add 3 or 4 grams of the fine sulphur. Shake the tube until all dis- solves, then pour the solution into a small beaker, which leave in a quiet place for spontaneous evaporation. The sulphur separates in octahedral crystals. Sulphur melts at a temperature of about 114 C. to a thin yellow liquid; at a higher temperature it grows darker and becomes viscid, so that it can be poured only with difficulty. This and other facts may be readily shown by trial. The boiling point is about 440. Sulphur crys- tallizes in several forms which exhibit different physical properties. A form which crystallizes in octahedra is found in nature. Similar rhombic octahedra are obtained by crys- tallization from carbon disulphide as explained above. These crystals have a specific gravity of 2.05 at 0. When 148 GENERAL CHEMISTRY. sulphur is melted and allowed to cool slowly, a portion of it separates in long needles of the monoclinic system. These have a lower specific gravity than the octahedral variety, viz. : 1.96. These monoclinic prismatic needles may be easily obtained by melting some sulphur in a test- tube and allowing it to cool a short time until a solid crust forms over the top. When this is broken through the still liquid portion may be poured out, leaving a crystalline mass attached to the sides of the tube. Ex. 105. Melt 15 to 20 Gm. of sulphur in a test-tube gradually, observe the changes in color and degree of fluidity. Above about 250 C. the melted mass grows thinner. After it has become quite thin pour it into some water and allow it to cool. A stringy plastic mass is obtained which is elastic like crude rubber. Set this aside and allow it to remain several days, and then observe that it has become brittle, or has returned to the common form. At the ordinary temperature the affinity of sulphur for the metals is slight, but at higher temperatures many com- binations may be easily made. It has been shown already that sulphur and copper unite very readily when strongly heated. The same reaction will now be shown with iron. Ex. 106. In a test-tube mix some flowers of sulphur with about an equal weight of finely divided iron. Fine filings or powder are prefer- able, drillings being usually too coarse. Heat the mixture slowly in the lamp flame. The sulphur melts and finally reaches a temperature at which chemical union between the two substances, accompanied by glowing of the mass, takes place. The iron burns with the sulphur as it does in oxygen and the product is sulphide of iron, or ferrous sul- phide. When the tube cools, break it, and examine the contents. Powder some in a mortar and observe the uniform dark color. Put some small pieces in a test-tube, add some water and then a little hydro- chloric acid. Observe that a gas with a very disagreeable odor is given off. Uses of Sulphur. Sulphur is employed in the manu- facture of sulphuric acid, in gunpowder (ordinarily a mix- ture of sulphur, saltpeter and charcoal), in many varieties of friction matches and in the preparation of several sul- phides. It is employed in considerable quantity in the vulcanization of rubber and in the preparation of so-called hard rubber. It has also numerous minor applications, in medicine and in the arts. GENERAL CHEMISTRY. 149 SULPHUR AND OXYGEN. Sulphur forms two important compounds with oxygen, one of which has been referred to before in the experiments on oxygen gas. At a temperature sufficiently elevated the union takes place directly as has been shown. The product formed is sulphurous oxide, SO 2 . Sulphurous Oxide. This oxide is found to a small extent in the atmosphere of cities where much soft coal is burned. It is also given off in some volcanic gases. For FIG. 17. experimental purposes it may be obtained by a peculiar reaction in which copper is made to decompose strong sul- phuric acid, as explained below. Ex. 107. Arrange apparatus as shown in the illustration. The flask should hold 400 to 500 Cc. Put in it 15 to 20 Gm. of copper in small pieces or turnings and add 25 Cc. of strong sulphuric acid. The Woulfe bottle contains a small amount of water to wash the gas passing through it, and the delivery tube from this can be led into another ves- sel of water, or into a clean dry bottle, for collection of the gas. Apply 150 GENERAL CHEMISTRY. heat to the flask and when gas bubbles begin to escape rapidly lower the flame to avoid too violent a reaction. Collect first several bottles of the gas, by displacement of air, which can be easily done as the gas is over twice as heavy as air. In collecting the gas cover the dry bottle as well as possible with a glass plate or perforated cardboard. When the bottle is full a burning match held at the mouth will be extinguished. As the bottles fill cover them and set aside for experiment. Then lead the delivery tube into a bottle of water and allow this to stand as long as gas is given off. The water absorbs the gas and is used below. For technical purposes sulphurous oxide may be made by heating charcoal with strong sulphuric acid. A reduc- tion of the acid takes place and oxide of carbon is formed as well as oxide of sulphur: C+2H 2 SO 4 =CO 2 4-2H 2 O-{-2SO 2 . For operations on the large scale, such as bleaching and fumigation, the oxide is made by burning sulphur in the air or by roasting certain sulphides called pyrites. Properties of Sulphurous Oxide. Of these, the odor is most characteristic, while in its chemical behavior sev- eral marked peculiarities may be easily shown. In col- lecting the gas the fact that it is not a supporter of com- bustion was shown. That sulphurous oxide is a good bleaching agent can be shown by experiments on printed cotton goods. At a low temperature the gas may readily be condensed to the liquid condition. The boiling point of the liquid is 8, and in this form it is an article of commerce, being employed in refrigeration. Ex. 108. Moisten a strip of calico and suspend it in a bottle of the gas, allowing it to remain half an hour. Many colors are com- pletely bleached in this time. It must be noted, however, that some colors are not at all acted upon by the gas. Ex. 109. The marked solubility of sulphurous oxide in water may be shown by means of a bottle well filled by the gas. Invert the bottle, closed by a glass plate, bring the mouth beneath the surface of water and then remove the plate. Water rushes up to take the place of the dissolved gas, as was the case in the experiment with ammonia. Lift up the bottle by means of the plate and observe that the water has an acid reaction, as shown by the litmus paper test. GENERAL CHEMISTRY. 151 At the temperature of 1 volume of water dissolves nearly 80 volumes of the gas; at 20 1 volume of water dis- solves 39.5 volumes of the gas; one liter of the gas weighs at 2.89 gm. Ex. HO. Moisten a long splinter of pine wood in strong nitric acid and dip it into a bottle of the gas. Red fumes are formed of oxides of nitrogen, showing the decomposition of the nitric acid. In this reaction the sulphurous oxide is said to act as a reducing agent. The meaning of this term will be explained below. We have remaining the solution of the gas in water, obtained by direct absorption. Tests may be made with it as follows: Ex. III. Pour portions of about 10 Cc. each into several test- tubes. To one add some solution of potassium permanganate, a few drops at a time. The deep purple color of this solution gives place to a very light pink. Into another portion pour a few cubic centimeters of a dilute solution of potassium dichromate. Notice the change of color to green. Boil a few cubic centimeters of ferric chloride in a test- tube and pour this, a few drops at a time, into the solution of sulphur- ous oxide, heating the latter after each addition. The yellowish brown color of the iron solution changes to pale green by this treatment. These are all reduction actions again. Ex. 112. Pour about 10 Cc. of the sulphurous oxide solution into a porcelain dish, which place on a sand-bath and heat. Everything evaporates showing that the product is volatile. The solution is called sulphurous acid, but, like the ammonium hydroxide, is not stable. In another dish pour about 25 Cc. of the solution, add to it some litmus, and then dilute solution of caustic soda until the color turns blue. Then restore the red color by addition of more sulphurpus oxide solution, put the dish on a sand-bath, and evaporate slowly to dryness. A white resi- due is left which has a sharp saline taste, quite distinct from that of the caustic soda. Now add a little water to the dish and some dilute hydro- chloric acid. Effervescence follows with escape of gas, which the odor shows is sulphurous xide. The caustic soda formed with the solution sodium sulphite, which is decomposed by the hydrochloric acid, with liberation of the sulphurous oxide and formation of sodium chloride. Ex. 113. Pour 10 Cc. of the sulphurous oxide solution into a test- tube and add to it a few drops of solution of barium chloride. A precipi- tate forms which consists of barium sulphite, and which dissolves readily by addition of a little hydrochloric acid. Now repeat the experiment, but add two or three drops of strong nitric acid to the sulphurous oxide solution, boil a few minutes and complete as before. A white precipi- tate forms here which does not dissolve on addition of hydrochloric acid. This is barium sulphate. 152 GENERAL CHEMISTRY. Leave the bottle containing the remainder of the sul- phurous oxide solution uncorked, but away from dust, for future experiments. We must now turn to a consideration of the chemical changes involved in the experiments above. First, we have the reaction by which the sulphurous oxide was produced. When copper dissolves in hot sulphuric acid copper sul- phate is formed, while water vapor and the sulphurous oxide are given off. Sulphuric acid contains, as already shown, hydrogen, oxygen and sulphur in the proportions given by the formula H 2 SO 4 . The liberation of the sul- phurous oxide, therefore, involves the breaking up of this group. It is possible that we have first a reaction analo- gous to one which was suggested as taking place between copper and nitric acid, viz.: Cu+H 2 SO 4 = CuSO 4 + H 2 , and that the liberated hydrogen acts here, as there, as a reducing agent; that is, one which decomposes because of its power of combining with oxygen. The second re- action, then, would be this : H 2 + H 2 SO 4 = 2H 2 O+ SO 2 Hydrogen + Sul a P c ^ ric = Water -+ If we consider these two reactions as taking place to- gether we may write : It appears from this last equation, which represents the results of experiments made quantitatively, that one- half of the sulphuric acid which takes part in the reaction becomes combined to form a sulphate, while the other half is broken up. yielding water and sulphurous oxide. Another explanation of the reaction between copper and sulphuric acid has been proposed and that is illus- trated by these equations : Cu-f-H 2 SO t =CuO + H 2 SO 3 , GENERAL CHEMISTRY. 153 copper oxide and sulphurous acid being formed. The first dissolves in the excess of sulphuric acid : 2 SO 4 =CuSO 4 and the second breaks up into water and sulphurous oxide: H 2 S0 3 =H 2 0+S0 2 . The experiments made with the sulphurous oxide in the condition of gas or in solution are mainly illustrative of one thing, viz., its reducing or oxygen absorbing power. In the experiment in which a splinter moistened with nitric acid was dipped into a jar of the gas it was observed that reddish fumes were given off. Now, it has been already explained that the appearance of these reddish fumes in reactions in which nitric acid is concerned is indicative of a decomposition of the acid by some substance which can take oxygen. The reddish fumes come from the residue left after this breaking up of the nitric acid. Sulphurous oxide, SO 2 , has a great tendency to take up oxygen and form a body called sulphuric oxide, SO 3 , and especially in presence of moisture. The nitric acid here furnishes the oxygen for the purpose. It was shown that the deep purple solution of potas- sium permanganate is completely decolorized by the solu- tion of sulphurous oxide. The reaction here is somewhat complicated, but is one in which the potassium perman- ganate parts with oxygen, converting the sulphurous oxide and water into sulphuric acid. The solution of the potas- sium dichromate, used in the same experiment, behaves quite in the same manner. The potassium dichromate acts here as an oxidizing agent. Because of their complexity these reactions will not be further explained here, but they are important ones and will be taken up later. We have next a case which is somewhat simpler. The solution of sulphurous oxide when treated with nitric acid forms sulphuric acid. We know this because after the mixing of the liquids we are able to prove the presence of 154 GENERAL CHEMISTRY. sulphuric acid by a reaction which will later be shown to be a certain test [or the acid. What the experiment, carried out in detail, shows to take place is illustrated by this equa- tion: O + 2HNO s =2H 8 S0 4 +N 8 3 . That is, an oxide of nitrogen and sulphuric acid are formed. In the last experiment made it was directed to allow a solution of sulphurous oxide in water to stand exposed to the air. This solution is often considered as one of sul- phurous acid, formed by direct union of the substances, thus : S0 2 -fH 2 = H 2 S0 3 . After standing some days or weeks exposed to the air we observe that a change has taken place in the liquid. We find that sulphuric acid is present, as we have tests by which we may readily distinguish sulphuric acid from sul- phurous. The latter acid takes up oxygen in this manner: Sulphuric acid is formed, therefore, as an oxidation product of sulphurous acid. This leads to the more de- tailed consideration of the properties of sulphuric acid, given later. SULPHUROUS ACID AND SULPHITES. As suggested above and as indicated by experiment, the solution of SO 2 in water may be considered as sul- phurous acid. As shown, the acid is not stable but decom- poses by heat and on exposure to the air. The sulphites are more stable and several are in common use. The fol- lowing equation illustrates their formation: ,SO,=Na SO s -h2H 2 O. When exposed to moist air most of the sulphites be- come sulphates by absorption of oxygen. GENERAL CHEMISTRY, 155 Uses of Sulphurous Oxide and Acid. It will be shown later that large quantities of sulphurous oxide are made as a step in the manufacture of sulphuric acid. It is also employed in bleaching silk, woolen and straw articles and in the fumigation of buildings. It is used also to pro- tect trees and vines from the ravages of certain pests. Strong solutions of sulphurous acid or of acid sulphites are employed in washing barrels and tanks or vats used in the manufacture and storage of beer. It acts here to de- stroy ferments, whose presence might spoil the product. In the laboratory a solution of sulphurous acid is often employed as a reagent. Hyposulphurous Acid. By the action of zinc on a solution of sulphurous oxide a peculiar acid is formed which is the true hyposulphurous acid. Zn+H 2 O+2SO 2 =:ZnSO 3 + H 2 SO 2 Zinc Hyposulphur- sulphite cms acid. The acid forms a yellow solution which absorbs oxygen readily and therefore acts as a strong reducing and bleach- ing agent. A corresponding salt is made by action of zinc on a solution of sodium acid sulphite: Zn+3NaHS0 8 =NaHS0 8 +Na 8 SO s +ZnS0 3 +H 8 0. The zinc and sodium sulphites may be crystallized out leaving a solution of the acid hyposulphite which is some- times used in bleaching. It has been used in the bleach- ing of syrups and other articles of food, but for this pur- pose its employment should be strongly condemned. Sulphuric Oxide or Sulphur Trioxide. This is a substance having the formula SO 3 , and is solid at the ordinary temperature. It may be made by several reac- tions, and readily when a mixture of sulphurous oxide and oxygen gas is passed over hot platinum sponge. This spongy material at a moderate temperature causes me two gases to combine. It may be more easily made by the dis- tillation of a liquid known as fuming sulphuric acid, which 156 GENERAL CHEMISTRY. will be described below. This acid has the composition H 2 S 2 O 7 , and when distilled decomposes forming H 2 SO 4 and SO 3 , which may be collected in a cool and perfectly dry receiver. Properties of Sulphur Trioxide. The substance made as described above appears in the form of dense, white fumes which yield long, silky needles on cooling. This crystalline solid must be preserved in sealed glass vessels. If brought in contact with the air it immediately attracts moisture and becomes liquid sulphuric acid: S0 3 +H 2 O^H 2 S0 4 . The trioxide itself melts at a temperature of 16 to form a colorless liquid. In dry form it is without action on litmus paper. SULPHURIC ACID. History. This very important acid has been known for many years, and was first made by the distillation of ef- floresced green vitriol, hence the name, oil of vitriol, which still clings to it. Reference to the acid appears first in the writings of the Arabian philosopher, Geber, and later, but very indefinitely, in the works of the earlier alchemists. In the fifteenth century Basil Valentine described more clearly the preparation of the vitriol and the distillation of the same. Practically all the acid used for 300 years was made by that reaction which will be referred to again, below. About the middle of last century a process was discovered by which sulphuric acid may be made by the oxidation of sulphurous oxide. This process was developed in Eng- land, while the other grew to importance in Germany. The common sulphuric acid used to day is made by a proc- ess which is a development of the crude attempts first made in England about 150 years ago. Preparation. The manufacture of sulphuric acid on the large scale involves a number of reactions illustrated GENERAL CHEMISTRY. 157 by what has been already given. First, sulphur, or a com- pound of iron and sulphur, known as iron pyrites, is burned in furnaces to form SO 2 by the aid of oxygen from the air. Then the gaseous sulphurous oxide, fumes of nitric acid and steam are led together into a large lead lined chamber in which they react on each other as illustrated by these equations: SO 2 + HNO 3 =HO(NO 2 )SO 2 . Nitroso-sulphuric acid. This first product, nitroso-sulphuric acid, is decom- posed by steam with formation of sulphuric acid and oxides of nitrogen: A fresh quantity of sulphurous oxide entering the lead chamber along with air drawn in and more steam combine to produce a new quantity of the nitroso-sulphuric acid. to be broken up by steam as before, making a continuous process. The nitroso product may also suffer decomposition in this manner: 2HO(NO 2 )SO 2 +SO 2 -f2H 2 O=:3H 2 SO 4 +2NO. The oxide of nitrogen, NO, is easily oxidized and on this the continuity of the process largely depends. It ap- pears, therefore, theoretically, that a very small amount of nitric acid is sufficient to convert an infinitely large amount of sulphurous acid into sulphuric acid. In practice this is not quite true as there is always some loss of the oxidizing gases. Oxygen is supplied by means of a current of air drawn into the chambers, and it is of course necessary to remove the residual nitrogen. In withdrawing this, small amounts of the nitrogen oxides escape and are lost to the process. Several secondary reactions take place which 158 GENERAL CHEMISTRY. also occasion slight losses. In the annexed illustration the general arrangement of the lead chambers and fur- naces are shown. Sulphur or pyrite is burned in stoves at A, and the FIG. 18. * fumes with excess of air pass up into the bottom of the tower, E. In the upper part of the sulphur burner a little nitric acid is generated by the action of a small amount of sulphuric acid on sodium nitrate and this nitric acid, in vapor, passes along with the air and sulphurous oxide. GENERAL CHEMISTRY. 159 Steam is generated in the boiler, B, and this is shown as entering the lead chambers at different points. As the several substances come together in the first lead cham- ber the reactions given above begin but they are not com- pleted until three or four are passed. In the third or largest chamber the principal part of the combination is completed. By time the last chamber is reached the sulphurous oxide is all in combination, but mixed with the nitrogen and oxygen from the air there is always some of the oxides of nitrogen. To save these the whole gaseous residue is drawn by means of a tall chimney or blower down into the bottom of the tower C, filled with coke or hard brick; a stream of strong sulphuric acid trickles down over this coke and absorbs the oxides, but allows the other gases to escape. This strong acid with its oxidizing ab- sorbed product flows finally into a reservoir, D, from which it is pumped to a tank above the first compartment, E, into which it is discharged slowly. Here it meets a stream of water. The dilution causes it to give up the dissolved oxides of nitrogen, because they are not soluble in weak sulphuric acid. The liberated oxides pass into the first lead chamber again with the sulphurous oxide from the burners and thus remain in circulation. To hasten oxida- tion a little fresh nitric acid is sometimes allowed to enter the second chamber and flow over the cascade there illus- trated. The coke tower to absorb the oxides of nitrogen is called a Gay Lussac tower, and the tower above the bur- ners where the strong mixture is discharged, is known as a Glover tower. The acid formed in the chambers has a specific gravity of about 1.6 and contains about 70 per cent of actual H 2 SO 4 . This is called chamber acid and is sold for many purposes. A more concentrated acid is produced by boiling down this chamber acid in leaden pans until a product of 1. 70 to 1.72 specific gravity is reached; this contains 77 to 80 per cent of acid and is strong enough for most chemical decomposi- tions. Concentration cannot be carried beyond this in lead because it dissolves in strong acid. The most con- centrated acid of commerce is made by evaporation in large glass globes or in platinum pans. The purest acid is 160 GENERAL CHEMISTRY. distilled from platinum stills, which increases its cost con- siderably. Properties. Pure sulphuric acid has a specific gravity of 1.85 at 0. It is colorless and oily in appearance. When heated it decomposes slightly, and if the temperature of Ebullition, about 338, is reached an acid of 98.5 per cent strength is left. This acid may be distilled without further change. Commercial sulphuric acid contains several im- purities and is often brownish in color from the presence of traces of organic matter. It is seldom stronger than 95 per cent, and is often below 90 per cent. The specific gravity of this acid is usually between 1.81 and 1.83 at 15. Large quantities of weaker acids are made and sold for special purposes. The so called chamber acid\s sold without con- centration, just as it leaves the lead chambers, and is used in several industries. Strong sulphuric acid has a remarkable action on water, which is shown in the next experiment : Ex. 114. Pour about 5 Cc. of water into a beaker or test-tube and add to it, stirring meanwhile, about double its volume of strong sul- phuric acid. Pour the acid into the ^vater, not the reverse. It will be no- ticed that the mixture becomes very hot, and that steam even may escape. This is a very characteristic behavior of the acid. Ex. 115. Pour 10 Cc., about, of strong sulphuric acid into a small beaker, which leave uncovered but protected from dust for several days. Notice at the start the depth to which the acid fills the beaker. From time to time look at it and observe that the liquid layer gradually grows deeper. If the beaker can be allowed to stand some weeks a very marked increase in the volume of the liquid will be seen. At the end of this long period repeat the last experiment by pouring the acid into water. A very high temperature on mixing will not be observed now. In fact, if the acid can be allowed to stand long enough, an increase in temperature may be scarcely perceptible. Another curious reaction depending on the same affinity of sulphuric acid for water is shown in the next experi- ment. In this case a compound is decomposed and the elements of water, hydrogen and oxygen, abstracted. Ex. 116. Make a very strong solution of cane sugar by dissolving about 20 Gm. in half its weight of water, in a beaker holding 300 Cc. or more. Pour into the syrup thus made an equal volume of strong sul- GENERAL CHEMISTRY. 161 phuric acid. In a few seconds the mixture becomes hot, blackens and gives off steam. A large volume of loose, finely divided carbon sepa- rates and rises to fill the beaker, carried up by the hot vapor. The sugar is a combination of carbon with hydrogen and oxygen; the last two are taken by the acid in the form of water, while the carbon is left in the free state. Ordinary pine wood has a composition very similar to that of sugar. When a splinter is dipped in strong sulphuric acid it blackens for the same reason. The best tests we have for the recognition of sulphuric acid in the free state depend on the principles just illus- trated. Sulphuric acid in combination, that is, in sul- phates, may be recognized by other tests, illustrated below. The behavior of strong sulphuric acid when mixed with salt, saltpeter and many other substances is also char- acteristic. It has been shown that in the case of salt a re- action follows in which hydrochloric acid and sodium sul- phate are formed. With saltpeter we obtain nitric acid and potassium sulphate, as shown. It can readily be proven that by heating a substance known as sodium ace- tate with strong sulphuric acid we obtain acetic acid and sodium sulphate. In general, it may be said, the mem- bers of a large and important class of bodies, known as salts, are decomposed by this acid, yielding sulphates and other acids. From chlorides we obtain hydrochloric acid; from nitrates, nitric acid; from acetates, acetic acid; from phosphates, phosphoric acid, and so on. More will be said about these salts later. We speak of sulphuric acid as a strong acid, because it is able to produce these decompo- sitions. The behavior of sulphuric acid with metals is inter- esting and has been illustrated by several experiments al- ready. Iron and-zinc, and many of the less common metals, are dissolved in sulphuric acid, especially when it is diluted, with formation of sulphate and liberation of hydrogen. Copper, mercury and some other metals are dissolved by sulphuric acid when it is concentrated and hot with formation of sulphates and liberation of sulphur- ous oxide instead of hydrogen, as shown in the case of copper. Uses of Sulphuric Acid. As has been shown, the manufacture of most of the other common acids depends on 162 GENERAL CHEMISTRY. the use of sulphuric acid. It is also largely used in the decomposition of common salt as a step in the manufacture of sodium carbonate by the Leblanc process. It is used in the decomposition of phosphate rocks in the manufac- ture of phosphatic fertilizers, which is a very important industry. In the manufacture of many organic dye-stuffs the use of strong sulphuric acid is necessary, and, in fact, it is used in some stage in hundreds of technical processes. The refining of petroleum, the production of glucose, the recovery of ammonia from gas liquor and the manufacture of most modern high explosives, are all industries in which this acid is practically necessary. It has become a com- monplace remark that the industrial development of a country may well be measured by the amount of sulphuric acid it uses. Oil of Vitriol or Pyrosulphuric Acid. It was said at the outset that sulphuric acid was first made by distilla- tion of green vitriol. If this substance, which has the formula- FeSO 4 .7H 2 O, is dried it loses most of its water and leaves a basic sulphate having the composition Fe 2 S 2 O 9 . When this residue is distilled it breaks up as here represented: Fe 2 S 2 9 =Fe 2 3 +2S0 3 . This SO 3 combines with the small amount of water left in the product and produces an acid of the composition H 2 S 2 O 7 . This is the real oil of vitriol, pyrosulphuric acid, or fuming sulphuric acid. It is made practically by dis- tilling the dried vitriol in small well glazed earthen retorts. When exposed to the air it gives off SO 3 , which combines with moisture to produce white fumes, H 2 SO 4 resulting. This very strong acid is made in large quanti- ties for use in the manufacture of organic products known as sulphonic acids. Some of these acids are bodies of great practical importance and certain ones among them cannot be made from the weaker acids. The name Nordhausen acid was formerly applied to this strong acid from the place in Germany where most of it was at one time made. GENERAL CHEMISTRY. 163 The Sulphates. Sulphates are salts formed from sulphuric acid by the replacement of its hydrogen by metals. Several illustrations of this have been given. They may be made, also, by combining sulphuric acid with many salts, as explained above. Many of these sulphates are important bodies and most of them are soluble in water. A few are not soluble, and on this fact a method for the recognition of the whole group is based. Ex. 117. Pour a weak solution of sodium sulphate into each of three test-tubes. To one add some solution of barium chloride, to the second some solution of strontium chloride, and to the third some solu- tion of lead acetate. In each case a white precipitate forms which, after a time, settles to the bottom of the test-tube; without waiting for this, however, add to each one of three tubes some hydrochloric acid and shake the mixture. No change is observed, then warm it, and still no change appears in the first two. A partial solution may result with the third. We have here, therefore, precipitates which are insoluble in water and hydrochloric acid and which contain barium, lead and stron- tium as sulphates. Of these the barium sulphate is the most character- istic and important. Barium chloride used as above is a test for sul- phates, and in works on analytical chemistry it is shown that when prop- erly employed it gives us accurate information concerning the presence of sulphates in complex mixtures even. The reactions illustrating these precipitations may be given here. First, with barium chloride and sodium sul- phate we have: BaCl 2 -fNa 2 SO 4 :=BaSO 4 -f-2NaCl Barium I Sodium _ Barium I Sodium chloride ' sulphate sulphate 'chloride. Then with strontium chloride we have: SrCl 2 +Na 2 SO 4 = SrSO 4 + 2NaCl Strontiuml Sodium - Strontium I Sodium chloride ' sulphate sulphate ' chloride. With lead acetate we have: Pb(C 2 H 3 2 ) 2 +Na 2 S0 4 = PbS0 4 +2NaC 2 H 3 2 Lead _!_ Sodium _ Lead _j_ Sodium acetate sulphate " sulphate ' acetate. Barium ' chloride solution with sulphuric acid itself gives the same white precipitate, but in this case hydro- chloric acid is formed: BaCl 8 +H 8 S0 4 =BaS0 4 +2HCl. 164 GENERAL CHEMISTRY. *| It is customary to speak oifree and combined sulphuric acid. The second expression, referring to the sulphates, is not strictly accurate as the whole of the acid is not in combination. But the characteristic part is and hence the use of the term. The same applies to other salts, as the nitrates, the chlorides, and others. All acids have hydro- gen in common, but with the hydrogen we have something characteristic for each acid. This characteristic element or group enters as combined sulphuric, nitric, hydrochloric or other acid into the corresponding salts. Other Sulphuric Acids. Besides the above several other acids containing hydro- gen, oxygen and sulphur are known, free or in combination. The most important one of these is called thiosulphuric acid. This exists in the well-known salt, sodium thiosul- phate, Na 2 S 2 O 3 -j-5H 2 O, formerly called hyposulphite of soda. This salt is largely used by photographers. The free acid is not stable. Solutions of thiosulphates de- compose with liberation of sulphurous oxide and pre- cipitation of sulphur when mixed with dilute acids. The names and formulas of other acids are these : Dithionic acid H S 2 O 6 . Trithionic acid H 2 S 3 O 6 . Tetrathionie acid H 2 S 4 O 6 . Pentathionic acid H 2 S 5 O 6 . These acids are rare and have no practical importance in the arts. A detailed description of them is not called for in an elementary book. Of somewhat greater importance are the persulphates and persulphuric acid, recently discovered. The acid has the formula HSO 4 , or H 2 S 2 O 8 , and is made by the elec- trolysis of strong H 2 SO 4 at a low temperature. This per- sulphuric acid is not stable, but decomposes with water yielding hydrogen peroxide : GENERAL CHEMISTRY. 165 It is therefore a strong oxidizing agent. Some of the persulphates are of technical importance. The potassium salt, KSO 4 or K 2 S 2 O 8 , is made by electrolysis of potas- sium acid sulphate, KHSO 4 at a low temperature. When warmed with water it decomposes in this way: It is employed as an oxidizer. An oxide correspond- ing has the formula S 2 O 7 , but is not technically useful in the pure state. It is not stable. SULPHUR AND HYDROGEN. Two combinations of these elements are known, but only one of them is important. This is the compound known as hydrogen sulphide. Hydrogen Sulphide, or sulphuretted hydrogen, occurs in the free state in some volcanic gases and certain mineral spring waters to which it imparts a marked odor. It is produced by the decomposition of vegetable matters containing sulphur, or from animal albuminoids. It can be formed by the direct union of the two elements, hydrogen and sulphur, at a high temperature, but is most easily made by a method quite analogous to that by which hydrochloric acid is made from a chloride, that is by the decomposition of a sulphide by means of some strong acid. We use for this purpose an artificial substance known as the sulphide of iron, or ferrous sul- phide, the preparation of which was illustrated in the experiment wherein iron and sulphur were melted together and strongly heated. When an acid is poured over the ferrous sulphide it is decomposed with production of the hydrogen sulphide, as illustrated by the following equation : FeS + H 2 S0 4 = H 2 S -f FeSO 4 Ferrous I Sulphuric =r Hydrogen I Ferrous sulphide ' acid sulphide 'sulphate. This method of making the hydrogen sulphide is shown in the following experiment: 166 GENERAL CHEMISTRY. Ex. 118. Fit a bottle with a funnel tube and delivery tube as shown in the figure below. Put in it about 20 to 25 Gm. of ferrous sul- phide in small lumps and pour through the funnel tube enough dilute sulphuric or hydrochloric acid to cover the sulphide and the end of the tube. The delivery tube leads down into a dry bottle covered as well as possible during the experiment with a glass plate. The gas is slightly heavier than air and can therefore be collected in this manner, but not perfectly. After adding the acid, a few minutes must be allowed for the expulsion of air by the liberated gas; then collect several bottles of the latter for experiment. With one bottle show the combustibility of the gas by burning it as with hydrogen ; observe the odor and test the reaction FIG. 19. of the product of the combustion by means of litmus paper. With a second bottle of the gas show its solubility in water. To this end cover the bottle with a glass plate, invert it with the mouth under water and then remove the plate. After a time water ascends into the bottle as the gas is absorbed. All this work must be done in a fume closet. The gas is somewhat poisonous. Ex. 119. After showing the properties of the gas, as above, re- place the dry bottles by one containing water. As the gas has been found to be to some degree soluble.a solution known as hydrogen sulphide water is thus obtained. While this solution is being made, prepare weak aqueous solutions of copper sulphate, lead acetate, and zinc ace- GENERAL CHEMISTRY. 167 tate and pour them into small flasks or beakers. Twenty-five Cc. of each solution will be sufficient. Remove the bottle containing the water from the generator and in its place put the flask or beaker with the copper sulphate. As the gas bubbles enter this solu- tion a black precipitate forms. In time this precipitate will settle to the bottom of the vessel. Now take off the delivery tube, wash and replace it and let the gas pass next into the solution of lead acetate. A black pre- cipitate forms here also. Next, after washing the tube, pass the gas into the zinc solution. A precipitate appears here, but it is white instead of black. After a few minutes add to the precipitating mixture half a dozen drops of hydrochloric acid and notice that the precipitate disap- pears or is dissolved by the acid. Then from a test-tube add ammonia water, a drop at a time, and observe that when a certain amount has been added the precipitate forms again. To the flasks containing the precipitates from the lead and copper solutions add a little hydrochloric acid and notice that the precipitates fail to disappear, as in the case of that from the zinc solution. After these tests have been made, allow the remainder of the gas to bubble into some ammonia water contained in a small flask. This yields a solution of ammonium sulphide, (NH 4 ) 2 S. Make all these experiments in the fume closet. Properties. In the foregoing, some of the most char- acteristic properties of the hydrogen sulphide have been shown. It dissolves in water just as hydrochloric acid does, but to a much less extent. Most of the gas can be easily expelled by boiling. One volume of water dissolves about two volumes of the gas at the ordinary temperature. Under pressure the gas can be condensed to a liquid which boils at 61. In the combustion of the gas we have a phenomenon reminding of the behavior of hydrogen, but the flame is weaker and the odor of the product of combustion charac- teristic. Sulphurous oxide is formed in this experiment, as is illustrated by the equation: H 2 S+30 = H 2 0+S0 2 . In the moist condition this gas gave the acid test with litmus paper. The most important behavior of the gas is shown, how- ever, by the precipitation of the three solutions. These precipitates are called sulphides and are formed by a double decomposition between the hydrogen sulphide and the 168 GENERAL CHEMISTRY. substances in solution. The following equations illustrate their production: CuSO 4 -f H 2 S = CuS + H 2 SO 4 Copper (Hydrogen __ Copper i Sulphuric sulphate I sulphide sulphide I acid Pb(C 2 H 3 O 2 ) 2 + H 2 S = PbS -f2HC 2 H 3 O 2 Lead (Hydrogen Lead I Acetic acetate sulphide sulphide ' acid Zn(C 2 H 3 O 2 ) 2 -f H 2 S = ZnS + 2HC 2 H 3 O 2 Zinc _|_ Hydrogen - Zinc Acetic acetate ' sulphide sulphide I acid. As experiment shows and as the equations illustrate, the solutions, though they may be neutral to begin with, must become acid during the precipitation, because acids are formed. The precipitates are therefore insoluble in these acids. That the sulphides of copper and lead are in- soluble in hydrochloric acid also was shown above, but zinc sulphide was found to be soluble in hydrochloric acid. This behavior is very suggestive; we are able by it to dis- tinguish between lead and zinc solutions, although both may be colorless. In the one case we get a black sulphide, insoluble in hydrochloric acid, while in the other we ob- tain a white precipitate soluble in hydrochloric acid. To more fully illustrate this important matter let the student carry out the following tests: Ex. 120. Prepare dilute solutions of mercuric chloride, antimony chloride, ferrous sulphate, zinc sulphate, calcium chloride and sodium chloride. Acidify each one with hydrochloric acid and pass in hydrogen sulphide gas from a generator as before. With the mercury and anti- mony solutions we obtain precipitates, in the first case black and in the other orange. But in the other cases no precipitates form, even after passing the gas a long time. While the gas is passing add to each solu- tion in turn some ammonia water, and note the result. With the first there is apparently no change, but with the second there is. The orange yellow precipitate dissolves to form a dark yellow solution. In the iron solution we have a black precipitate and in the zinc a white, while in the other solutions no precipitate appears. Add now to these last some solution of ammonium carbonate. In the calcium solution a white pre- cipitate is formed while the other remains clear. It appears from the above that the mercury compound yields a precipitate in the presence of both acid and alkali, GENERAL CHEMISTRY. 169 the antimony compound 'gives a precipitate from the acid solution only, the iron and zinc compounds from the alka- line solution only, while the calcium and sodium com- pounds give no sulphide precipitates. The iron precipitate is distinguished from the zinc precipitate by its color, while finally the calcium and sodium compounds which give no precipitates with the gas are distinguished from each other by the fact that one yields a precipitate with am- monium carbonate, while the other does not. It should be added here that zinc may be precipitated in presence of acetic acid, as well as from alkaline solutions. These experiments are of fundamental importance, and the student will learn later that they are of common appli- cation in the branch of chemistry known as qualitative analysis. The hydrogen sulphide is one of our most important test substances and by its aid we are able not only to recognize bodies in solutions, but to make separa- tions 1 of bodies into groups, and thus isolate them from each other. We have remaining our solutions of hydrogen sulphide in water and in ammonia water. With these, experiments may be made to show that they behave in many cases as does the gas. Let the student determine for himself their action with solutions. Both are important reagents in the laboratory. SULPHUR AND CHLORINE. Three combinations of sulphur with chlorine are known. Sulphur Monochloride, S 2 C1 2 . This is a yellowish brown liquid easily made by passing dry chlorine over sul- phur melted in a retort. The liquid has a specific gravity of 1.705 and boils at 138. It dissolves sulphur readily and in large quantity. The solution so made is used in vul- canizing rubber. Sulphur Dichloride, SC1 2 , is made by passing dry chlorine into the monochloride at a temperature kept near the zero point. It is a dark liquid which decomposes if 170 GENERAL CHEMISTRY heated to about 20 or above, yielding free chlorine and the monochloride. Sulphur Tetrachloride, SC1 4 , is made by saturating the dichloride with chlorine at a temperature of 22. It is a light yellow liquid which decomposes quickly at tem- peratures above 22 Sulphur forms also compounds with iodine and bro- mine, but they are not important. SELENIUM. This is a comparatively rare element which resembles sulphur in many respects. It is a solid substance with a dark gray color and melts at about 217. The compounds of selenium resemble those of sul- phur. The best known are: selenium dioxide, SeO 2 , a white crystalline substance which dissolves in water to form selenous acid, H 2 SeO 3 ; selenic acid, H 2 SeO 4 , made by oxi- dation of a selenite; hydrogen selenide, H 2 Se, a gas resem- bling H 2 S and made by decomposing a selenide by an acid, and which precipitates many metals as does the sulphide. Selenium burns with a disagreeable odor described as resembling that of rotten cabbage, forming an oxide of un- known composition. TELLURIUM. Tellurium is another rare element of the sulphur group, found usually in combination as a telluride. In appear- ance it resembles the metals, but behaves chemically as do sulphur and selenium. Its specific gravity is 6.2 and its melting point is about 500. Its most important com- pounds are: hydrogen telluride, H 2 Te, a gas; tellurium di- oxide, TeO 2 , a white solid which forms tellurous acid, H 2 TeO 3 , with water; telluric acid, H 2 TeO 4 , a white solid, soluble in water. The three elements just considered constitute an inter- esting natural group in which the properties of the ele- ments themselves and of their compounds are functions of the corresponding atomic weights. This is shown in the following table : GENERAL CHEMISTRY. 171 Sulphur. Selenium. Tellurium. Atomic weight Specific gravity. 32.1 2.05 79.0 4.6 127.5 6.2 Melting point 114 217 500 Boiling point 440 665 Above 1,000 Hydrogen compound ous oxide H 2 S, gas. SO 2 a gas be- H 2 Se, gas. SeO 2 , a solid H 2 Te, gas. TeO 2 ,a crys- ous acid comes liquid at 8, soluble in water. H 2 SO 3 not readily soluble in water. talline solid, slightly soluble in water. H 2 TeO 3 , a ic oxide stable in free state. SO 3 volatile Not known. solid. TeO 8 , yellow ic acid solid. H 2 SO 4 liquid H 2 SeO 4 , crystalline solid. H 2 TeO 4 , white volatile acid not decomposed by HC1. heavy colorless liquid, decom- posed by HC1. solid mass, with water H 2 TeO 4 -f2H 2 O, crys- talline solid. It will be noticed that the specific gravities, the melt- ing points and the boiling points of the elements increase with the atomic weights. Also that the compounds become heavier or more nearly solids in the same order. It will be pointed out later that similar relations exist between the members of other groups and that in a general way the properties of elements are closely dependent on their atomic weights. CHAPTER VIII. SILICON AND BORON AND THEIR COMPOUNDS. THESE elements occur in nature combined with oxygen or with oxygen and metals. SILICON. This is one of the very abundant elements in combina- tion and is found as the oxide, SiO 2 , in several minerals of which the most common is quartz. Other substances, flint, white sand, opal, chalcedony and agate, consist essen- tially of this oxide. In many silicates the element is widely distributed, and it follows oxygen in point of abun- dance in the earth's crust. Preparation. The element may be separated by decomposing one of its compounds by potassium, by aid of heat: It is left after this reaction as an amorphous powder. If this is melted with zinc it becomes crystalline as the zinc cools and may be secured in this form by dissolving the metal. Its specific gravity is 2.49 and it is hard enough to scratch glass. As the methods followed in the isolation of silicon are expensive the free element has no technical uses. SILICON AND OXYGEN. But one oxide of silicon is known and this has the com- position SiO 2 , and is called silica. It occurs in very pure form in varieties of quartz and tridymite, both characteris- GENERAL CHEMISTRY. 173 tic crystalline minerals. In the opal it is found in amor- phous condition. Preparation. Pure silica may be easily made by de- composing a solution of sodium silicate, known as soluble glass, by means of hydrochloric acid. The precipitate which forms is thoroughly washed with water, dried and ignited. This leaves the silica in a fine amorphous condi- tion. Properties. Silica is practically insoluble in cold water and common acids. It is dissolved, however, by hydrofluoric acid, on which fact the etching of glass de- pends. At a very high temperature water (superheated) dissolves silica to a slight extent, forming silicic acid. In alkali solutions, especially if warm, amorphous silica dis- solves readily to form silicates. The crystalline varieties of silica may be converted into the same silicates by fusion with alkalies. SILICIC ACIDS. Silicon has a valency of four and the acid correspond- ing to it with the greatest molecular weight has the for- mula H 4 O 4 Si, and is known as orthosilicic acid. Its com- position is represented in this manner: H O\ H O | c . H-0/ Sl H O 7 This acid exists in solution, but as it is not stable, can- not be obtained in the free state. An acid derived from this is known: H 4 SiO 4 H 2 O = H 2 SiO 3 , H O x >Si = O. H O 7 To prepare solutions of these acids a weak solution of water-glass is decomposed by hydrochloric acid, leaving 174 GENERAL CHEMISTRY. orthosilicic acid dissolved. The mixture is thrown on a dialyzer, floating on water, and allowed to remain until it is free from hydrochloric acid and chlorides. These sub- stances pass through the membrane bottom of the dialyzer, but the colloidal silicic acid cannot. In this manner it is possible to prepare a weak, pure solution of the ortho acid. This may be concentrated to a strength of about 14 per cent. When evaporated beyond this, water is lost and the gelatinous acid, H 2 SiO 3 , is formed. This in turn by loss of water becomes SiO 2 . H 2 Si0 3 H 2 0: Silicates. Corresponding to the silicic acids a large number of bodies called silicates are known. Some of these can be formed artificially, but most of them occur in nature as mineral species, many of which are common and impor tant bodies. The composition of most of these minerals appears quite complex, but a little study shows their rela- tion to orthosilicic acid. For instance, the mineral serpen- tine may be represented by the formula Mg 3 Si 2 O 7 , which corresponds to an acid, H 6 Si 2 O 7 . Now, this in turn is related to the ortho acid, as illustrated: 2H 4 SiO 4 = H 6 Si 8 7 +H 8 0. The common mineral, orthoclase, is called a trisilicate, and is represented essentially by the formula AlKSi 3 O 8 , corresponding to H 4 Si 3 O 8 . This is 3H 4 Si0 4 4H 2 = H 4 Si 3 8 . It appears, therefore, that these silicates may be looked upon as derived from condensed silicic acids, formed from orthosilicic acid by loss of water. It will be pointed out that boric acid behaves much in the same way. The silicates of the alkali metals are soluble in water. GENERAL CHEMISTRY. 175 Potassium silicate and sodium silicate are called soluble glass or water-glass. Ex. 121. Take about 5 Cc. of the strong solution of sodium sili- cate, known as water-glass, and add to it, a little at a time, some con- centrated hydrochloric acid. When the mixture becomes quite strongly acid a gelatinous mass is produced, which becomes so stiff that the test-tube in which it is formed may be inverted without spilling it. The thick, colloidal substance is impure orthosilicic acid and metasilicic acid. If, before adding the hydrochloric acid, the water-glass is largely diluted with water no separation of the colloidal substance takes place. It remains in solution and can be partially purified by dialysis, as explained above. Silicic acid forms insoluble salts with many basic bod- ies, and some of these can be made by precipitation, as shown below : Ex. 122. Dilute the common water-glass with about 20 parts of water. Take small portions of this diluted liquid in test-tubes and add to them solutions of calcium chloride, copper sulphate, lead nitrate and cobalt nitrate. Precipitates are formed which are silicates of the metals in these salts. The soluble glass has approximately the composition Na 2 SiO 3 and the insoluble silicates may be made from it by double decomposition, as : Na 2 SiO 3 +CaCl 2 =:CaSiO 3 +2NaCl Sodium (Calcium Calcium _|_ Sodium silicate 'chloride silicate ' chloride. A soluble salt is left in the liquid. Advantage is taken of this behavior of the soluble silicate, or water-glass, in making certain kinds of cement and artificial stone. As mentioned, the alkali silicates are soluble; the others are insoluble in water and many of them cannot be decomposed by acids. Common glass is an artificial mix- ture of silicates made by fusing quartz sand, silica, with basic substances. For example, common window glass is made by fusing a mixture, in certain proportions, of sand, lime or pure limestone and dry sodium carbonate. Roughly speaking, we distinguish four varieties of 176 GENERAL CHEMISTRY. glass, viz. : crown or window glass, Bohemian glass, flint glass and common bottle glass. Crown Glass is essentially a mixture of calcium and sodium silicates. In some kinds a little alumina is present. It is made by melting at a very high heat a mixture of white sand, lime or limestone, and soda ash or dry sodium sulphate. Common window and plate glass and much hollow ware are included under crown glass. Bohemian Glass consists essentially of the silicates of potassium and calcium. It is made of carefully selected materials, usually quartz sand, pure refined potassium car- bonate and chalk, or well burned lime, as free as possible from magnesia. This glass can be fused only at a high temperature, and softens only with difficulty when heated. It is, therefore, employed in making much chemical glass- ware. Sometimes a little sodium carbonate is used with the potassium carbonate to make it more readily workable. Flint Glass is essentially a lead potassium silicate and is made by melting a combination of sand, red lead and dry potassium carbonate. This glass can be melted and cast or otherwise worked with comparative ease, and is therefore employed in making tableware and large arti- cles of ornamentation. The ready fusibility depends on the presence of lead silicate. This glass cannot be used for chemical ware. Common Green Bottle Glass resembles crown glass, but is made of impure materials. It usually contains con- siderable quantities of iron and aluminum silicates. The green color is due to the ferrous salt. The chemical composition of several kinds of glass as found by analysis is given in the following table, the re- sults being expressed in the usual manner. It must be remembered, however, that certain special kinds of glass contain still other substances. A part of the silicic acid may be replaced by boric acid and for some purposes the oxide of lead may be partly replaced by oxide of thallium. The relation of glass to pottery will be shown later. GENERAL CHEMISTRY. 177 KIND OF GLASS. Si0 2 Na 2 K 2 CaO MgO PbO Fe 2 3 A1 2 3 BOHEMIAN. Combustion tubing 74.19 76.41 1.87 1.38 13.13 10.96 9.39 9.71 0.36 0.49 89 Optical glass 75 81 2 00 15 03 12 13 32 1 02 Mirror plate 75.81 4.84 1L39 7.38 0.10 1,01 FLINT GLASS. German 75.24 12.51 1.48 10.48 English 51 40 9 40 37 40 2 00 Optical 44.30 11.75 43.05 0.50 0,12 CROWN GLASS. German window 71.56 12.97 13.27 1. 29 English window. 70 71 13 25 13 38 1. 02 French plate 73.00 11.50 15.50 English plate . 77 90 12 53 1.72 4.85 3. 59 German plate 78.75 13.00 6.50 1. 75 Coloring Glass. Certain metallic oxides may be melted with the glass mixture and so impart to the finished glass some desired shade. The red oxide of copper is used in making ruby glass, while the black oxide of the same metal gives a green color; ferrous oxide yields a green glass and ferric oxide a yellowish brown; the black oxide of manganese is used in giving a pink to purple, oxide of cobalt a deep blue, oxide of uranium a beautiful canary yellow. Various shades may be made by properly com- bining some of these oxides, and it is also possible by the proper combination to secure from impure materials an almost colorless glass. Black oxide of manganese is com- monly employed to correct the objectionable color due to presence of iron. This it does by oxidizing the iron to the ferric condition, the yellow tint of which is complementary to the purple of the manganese compound. SILICON AND HYDROGEN. One compound of these two substances is known, hav- ing the composition SiH 4 . It is a gaseous body, made by the action of acids on magnesium silicide and is an inter- esting compound from a theoretical standpoint, but has no technical applications. 178 GENERAL CHEMISTRY. SILICON AND THE HALOGENS. Silicon exists in combination with fluorine, chlorine, bromine and iodine. Of these the tetrafluoride, SiF 4 , is the most important. The formation of this in the etching of glass has been referred to already. It is produced by the action of hydrofluoric acid on silica: 4HF+Si0 2 =SiF 4 +2H 2 0. It is a gaseous substance which is decomposed by con- tact with water. The compounds SiHCl 3 and SiCl 4 are known. They are volatile liquids which decompose when mixed with water. Fluosilicic Acid. When silicon tetrafluoride is passed into water it decomposes in this way: 2SiF 4 +3H 2 O = H 2 SiF 6 -f2HF+H 2 SiO 3 . The body, H 2 SiF 6 , is known as fluosilicic acid. It is stable only in solution, and in this form is sometimes used as a reagent. BORON. This is an element which is found in a few natural sub- stances, of which borax, boric acid and calcium borate are the most important. The element may be liberated by the decomposition of some of its compounds, but it is not im- portant in the free state. The specific gravity of crystal- lized boron is 2.68. BORON AND OXYGEN. One compound is known having the composition B 2 O 3 . It is a glass like body, soluble in water, made best by strongly heating boric acid. BORIC ACID. This is a combination of boron with hydrogen and oxy- gen, having the composition H 3 BO 3 . It is found in na- GENERAL CHEMISTRY. 179 ture in small amount, and especially in the vapor from vol- canic fissures existing in certain parts of Tuscany. The water condensed from this vapor is collected in small lagoons, and kept boiling by the action of the hot vapor itself. In this way a rapid concentration is effected. Much of the boric acid of commerce comes from this source. It can be made from borax, however, and this will be illustrated here. Ex. 123. By the aid of heat dissolve about 30 Gm. of powdered borax in about 120 Cc. of water. Add to tne hot solution enough strong hydrochloric acid to make the liquid strongly acid to litmus paper. Stir well while adding the acid. Then allow the mixture to cool thoroughly. Thin crystalline plates of boric acid separate. Remove the supernatant liquid by filtration, take up the boric acid with hot water and purify it by recrystallization. Ex. 124. Boric acid is much more soluble in hot water than in cold. It is also readily soluble in alcohol. Prove this by dissolving the product of the last experiment in some alcohol. Pour some of the solu- tion so obtained over a little asbestos in a porcelain dish. When the asbestos is thoroughly moistened take it up with clean forceps and hold it in the flame of the Bunsen burner to ignite the alcohol. The flame produced is colored an intense green by the hot vaporized boric acid. This is a characteristic reaction and is employed for the recognition of boron compounds. Considerable quantities of boric acid and borax are now made from borocalcite, CaB 4 O 7 .4H 2 O, which occurs abundantly in California. Borax. This substance is a salt containing boric acid and sodium. Its chemical composition is shown by the formula Na 2 B 4 O 7 -f-10H 2 O. The water here represented is known as water of crystallization, and can be separated by heat, leaving what is known as borax glass or anhydrous borax. The preparation of this can be shown by a simple test. Ex. 125. Bend the end of a piece of platinum wire so as to form a loop two or three millimeters across. Heat this in the flame of the Bun- sen burner and then dip it while hot in some powdered borax. Heat again and repeat the operation until the loop is well covered. Then hold this in the flame several minutes. The mass swells and gives off steam, but finally fuses together and forms a colorless, clear, glass-like globule, filling the loop, and called the borax bea. Iodine .. 126 85 Cobalt .... 58. 93 | Platinum . . ...194 89 The properties of the elements in these groups vary with changes in the atomic weights. In the iron group the weights are nearly the same and we find that the metals and their compounds are much alike. The same is true in the platinum group. A consideration of such rela- tions would seem to suggest that the properties of ele- ments may depend on their atomic weights, and this has been shown to be in a marked degree the case. THE PROPERTIES OF THE ELEMENTS AS PERIODIC FUNCTIONS OF THEIR ATOMIC WEIGHTS. In order to show any relations existing between the atomic weights and properties of elements let us write their symbols in the order of increasing weights, beginning with lithium, as follows: Li, 7.03; Be, 9.08; B, 10.95; C, 12.01; N, 14.04; O, 16.00; F, 19.08. We have here a gradual increase in the atomic weights and a well characterized change in properties correspond- ing. Lithium is strongly metallic and positive in its behavior, while fluorine is as certainly nonmetallic and negative. The next greatest atomic weight is that of sodium = 23. 05. But we have here an element with prop- erties like those of lithium rather than like those of fluorine. Sodium evidently does not follow the latter ele- GENERAL CHEMISTRY. 251 ment in the series as begun. Let us therefore make a new series, parallel with the first, which runs thus: Na, 23.05; Mg, 24.28; Al, 27.11; Si, 28.40; P, 31.02; S, 32.07; Cl, 35.45. We have a new series of seven elements, beginning with a characteristic metal and ending with a characteristic nonmetal. The corresponding elements in the two series resemble each other very closely in their chemical behav- ior and in the compounds they form. It appears from this that certain properties are repeated in passing through a series or period of seven elements, involving a change in atomic weight of about 16 units. What has been done here for fourteen elements may be done for many more. This grouping was suggested nearly 30 years ago by Lothar Meyer and D. Mendelejeff, independently, and is called the Periodic Arrangement of the elements. It is shown in the following table in which the symbols and atomic weights are given, and also the differences in atomic weights in passing from one period to another. Some of the elements cannot be well included in the seven families or groups as first indicated and are therefore placed in a separate or eighth group. This periodic arrangement of the elements is often called the natural arrangement, but the appropriateness of this term may not be immediately apparent. The close relations of the elements of the first group, Li, Na, K, Rb and Cs and their compounds are easily recognized, but the position of the other elements, Cu, Ag and Au is not as evident. In studying the compounds of the metals later the student will find that the strongest likenesses are found between these compounds rather than between the metals themselves. In the next group the metals Be, Mg, Ca, Sr and Ba are very closely related, while among their com- pounds there are the closest resemblances. The salts of Zn and Cd are in many instances much like those of Mg, and hence the propriety of grouping them all together. The evidence for the position of Hg is not as clear. 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