AN ELEMENTARY STUDY OF CHEMISTRY BY WILLIAM McPHERSON, PH. D. PROFESSOR OF CHEMISTRY, OHIO STATE UNIVERSITY AND WILLIAM EDWARDS HENDERSON, PH. D. ASSOCIATE PROFESSOR OF CHEMISTRY, OHIO STATE UNIVERSITY _REVISED EDITION_ GINN & COMPANYBOSTON * NEW YORK * CHICAGO * LONDON COPYRIGHT, 1905, 1906, BYWILLIAM MCPHERSON AND WILLIAM E. HENDERSON ALL RIGHTS RESERVED The Athenęum PressGINN & COMPANY * PROPRIETORS * BOSTON * U. S. A. Transcriber's note: For Text: A word surrounded by a cedilla such as ~this~ signifies thatthe word is bolded in the text. A word surrounded by underscores like_this_ signifies the word is italics in the text. The italic and boldmarkup for single italized letters (such as variables in equations) and"foreign" abbreviations are deleted for easier reading. For numbers and equations: Parentheses have been added to clarifyfractions. Underscores before bracketed numbers in equations denote asubscript. Superscripts are designated with a caret and brackets, e. G. 11. 1^{3} is 11. 1 to the third power. Appendix A and B have been moved to the end of the book. Minor typos have been corrected. PREFACE In offering this book to teachers of elementary chemistry the authorslay no claim to any great originality. It has been their aim to preparea text-book constructed along lines which have become recognized as bestsuited to an elementary treatment of the subject. At the same time theyhave made a consistent effort to make the text clear in outline, simplein style and language, conservatively modern in point of view, andthoroughly teachable. The question as to what shall be included in an elementary text onchemistry is perhaps the most perplexing one which an author mustanswer. While an enthusiastic chemist with a broad understanding of thescience is very apt to go beyond the capacity of the elementary student, the authors of this text, after an experience of many years, cannot helpbelieving that the tendency has been rather in the other direction. Inmany texts no mention at all is made of fundamental laws of chemicalaction because their complete presentation is quite beyond thecomprehension of the student, whereas in many cases it is possible topresent the essential features of these laws in a way that will be ofreal assistance in the understanding of the science. For example, it isa difficult matter to deduce the law of mass action in any very simpleway; yet the elementary student can readily comprehend that reactionsare reversible, and that the point of equilibrium depends upon, rathersimple conditions. The authors believe that it is worth while topresent such principles in even an elementary and partial manner becausethey are of great assistance to the general student, and because theymake a foundation upon which the student who continues his studies tomore advanced courses can securely build. The authors have no apologies to make for the extent to which they havemade use of the theory of electrolytic dissociation. It is inevitablethat in any rapidly developing science there will be differences ofopinion in regard to the value of certain theories. There can be noquestion, however, that the outline of the theory of dissociation herepresented is in accord with the views of the very great majority of thechemists of the present time. Moreover, its introduction to the extentto which the authors have presented it simplifies rather than increasesthe difficulties with which the development of the principles of thescience is attended. The oxygen standard for atomic weights has been adopted throughout thetext. The International Committee, to which is assigned the duty ofyearly reporting a revised list of the atomic weights of the elements, has adopted this standard for their report, and there is no longer anyauthority for the older hydrogen standard. The authors do not believethat the adoption of the oxygen standard introduces any realdifficulties in making perfectly clear the methods by which atomicweights are calculated. The problems appended to the various chapters have been chosen with aview not only of fixing the principles developed in the text in the mindof the student, but also of enabling him to answer such questions asarise in his laboratory work. They are, therefore, more or lesspractical in character. It is not necessary that all of them should besolved, though with few exceptions the lists are not long. The answersto the questions are not directly given in the text as a rule, but canbe inferred from the statements made. They therefore require independentthought on the part of the student. With very few exceptions only such experiments are included in the textas cannot be easily carried out by the student. It is expected thatthese will be performed by the teacher at the lecture table. Directionsfor laboratory work by the student are published in a separate volume. While the authors believe that the most important function of theelementary text is to develop the principles of the science, theyrecognize the importance of some discussion of the practical applicationof these principles to our everyday life. Considerable space istherefore devoted to this phase of chemistry. The teacher shouldsupplement this discussion whenever possible by having the class visitdifferent factories where chemical processes are employed. Although this text is now for the first time offered to teachers ofelementary chemistry, it has nevertheless been used by a number ofteachers during the past three years. The present edition has beenlargely rewritten in the light of the criticisms offered, and we desireto express our thanks to the many teachers who have helped us in thisrespect, especially to Dr. William Lloyd Evans of this laboratory, ateacher of wide experience, for his continued interest and helpfulness. We also very cordially solicit correspondence with teachers who may finddifficulties or inaccuracies in the text. The authors wish to make acknowledgments for the photographs andengravings of eminent chemists from which the cuts included in the textwere taken; to Messrs. Elliott and Fry, London, England, for that ofRamsay; to The Macmillan Company for those of Davy and Dalton, takenfrom the Century Science Series; to the L. E. Knott Apparatus Company, Boston, for that of Bunsen. THE AUTHORS OHIO STATE UNIVERSITY COLUMBUS, OHIO CONTENTS CHAPTER PAGE I. INTRODUCTION 1 II. OXYGEN 13 III. HYDROGEN 28 IV. WATER AND HYDROGEN DIOXIDE 40 V. THE ATOMIC THEORY 59 VI. CHEMICAL EQUATIONS AND CALCULATIONS 68 VII. NITROGEN AND THE RARE ELEMENTS IN THE ATMOSPHERE 78 VIII. THE ATMOSPHERE 83 IX. SOLUTIONS 94 X. ACIDS, BASES, AND SALTS; NEUTRALIZATION 106 XI. VALENCE 116 XII. COMPOUNDS OF NITROGEN 122 XIII. REVERSIBLE REACTIONS AND CHEMICAL EQUILIBRIUM 137 XIV. SULPHUR AND ITS COMPOUNDS 143 XV. PERIODIC LAW 165 XVI. THE CHLORINE FAMILY 174 XVII. CARBON AND SOME OF ITS SIMPLER COMPOUNDS 196 XVIII. FLAMES, --ILLUMINANTS 213 XIX. MOLECULAR WEIGHTS, ATOMIC WEIGHTS, FORMULAS 223 XX. THE PHOSPHORUS FAMILY 238 XXI. SILICON, TITANIUM, BORON 257 XXII. THE METALS 267 XXIII. THE ALKALI METALS 274 XXIV. THE ALKALINE-EARTH FAMILY 300 XXV. THE MAGNESIUM FAMILY 316 XXVI. THE ALUMINIUM FAMILY 327 XXVII. THE IRON FAMILY 338 XXVIII. COPPER, MERCURY, AND SILVER 356 XXIX. TIN AND LEAD 370 XXX. MANGANESE AND CHROMIUM 379 XXXI. GOLD AND THE PLATINUM FAMILY 390 XXXII. SOME SIMPLE ORGANIC COMPOUNDS 397 INDEX 421 APPENDIX A Facing back cover APPENDIX B Inside back cover LIST OF FULL-PAGE ILLUSTRATIONS PAGEANTOINE LAURENT LAVOISIER _Frontispiece_ JOSEPH PRIESTLEY 14 JOHN DALTON 60 WILLIAM RAMSAY 82 DMITRI IVANOVITCH MENDELÉEFF 166 HENRI MOISSAN 176 SIR HUMPHRY DAVY 276 ROBERT WILHELM BUNSEN 298 AN ELEMENTARY STUDY OF CHEMISTRY CHAPTER I INTRODUCTION ~The natural sciences. ~ Before we advance very far in the study of nature, it becomes evident that the one large study must be divided into anumber of more limited ones for the convenience of the investigator aswell as of the student. These more limited studies are called the_natural sciences_. Since the study of nature is divided in this way for mere convenience, and not because there is any division in nature itself, it often happensthat the different sciences are very intimately related, and a thoroughknowledge of any one of them involves a considerable acquaintance withseveral others. Thus the botanist must know something about animals aswell as about plants; the student of human physiology must knowsomething about physics as well as about the parts of the body. ~Intimate relation of chemistry and physics. ~ Physics and chemistry aretwo sciences related in this close way, and it is not easy to make aprecise distinction between them. In a general way it may be said thatthey are both concerned with inanimate matter rather than with living, and more particularly with the changes which such matter may be made toundergo. These changes must be considered more closely before adefinition of the two sciences can be given. ~Physical changes. ~ One class of changes is not accompanied by analteration in the composition of matter. When a lump of coal is brokenthe pieces do not differ from the original lump save in size. A rod ofiron may be broken into pieces; it may be magnetized; it may be heateduntil it glows; it may be melted. In none of these changes has thecomposition of the iron been affected. The pieces of iron, themagnetized iron, the glowing iron, the melted iron, are just as trulyiron as was the original rod. Sugar may be dissolved in water, butneither the sugar nor the water is changed in composition. The resultingliquid has the sweet taste of sugar; moreover the water may beevaporated by heating and the sugar recovered unchanged. Such changesare called _physical changes_. DEFINITION: _Physical changes are those which do not involve a change inthe composition of the matter. _ ~Chemical changes. ~ Matter may undergo other changes in which itscomposition is altered. When a lump of coal is burned ashes andinvisible gases are formed which are entirely different in compositionand properties from the original coal. A rod of iron when exposed tomoist air is gradually changed into rust, which is entirely differentfrom the original iron. When sugar is heated a black substance is formedwhich is neither sweet nor soluble in water. Such changes are evidentlyquite different from the physical changes just described, for in themnew substances are formed in place of the ones undergoing change. Changes of this kind are called _chemical changes_. DEFINITION: _Chemical changes are those which involve a change in thecomposition of the matter. _ ~How to distinguish between physical and chemical changes. ~ It is notalways easy to tell to which class a given change belongs, and manycases will require careful thought on the part of the student. The testquestion in all cases is, Has the composition of the substance beenchanged? Usually this can be answered by a study of the properties ofthe substance before and after the change, since a change in compositionis attended by a change in properties. In some cases, however, only atrained observer can decide the question. ~Changes in physical state. ~ One class of physical changes should be notedwith especial care, since it is likely to prove misleading. It is afamiliar fact that ice is changed into water, and water into steam, byheating. Here we have three different substances, --the solid ice, theliquid water, and the gaseous steam, --the properties of which differwidely. The chemist can readily show, however, that these three bodieshave exactly the same composition, being composed of the same substancesin the same proportion. Hence the change from one of these substancesinto another is a physical change. Many other substances may, undersuitable conditions, be changed from solids into liquids, or fromliquids into gases, without change in composition. Thus butter and waxwill melt when heated; alcohol and gasoline will evaporate when exposedto the air. _The three states--solid, liquid, and gas--are called thethree physical states of matter. _ ~Physical and chemical properties. ~ Many properties of a substance can benoted without causing the substance to undergo chemical change, and aretherefore called its _physical properties_. Among these are its physicalstate, color, odor, taste, size, shape, weight. Other properties areonly discovered when the substance undergoes chemical change. These arecalled its _chemical properties_. Thus we find that coal burns in air, gunpowder explodes when ignited, milk sours when exposed to air. ~Definition of physics and chemistry. ~ It is now possible to make ageneral distinction between physics and chemistry. DEFINITION: _Physics is the science which deals with those changes inmatter which do not involve a change in composition. _ DEFINITION: _Chemistry is the science which deals with those changes inmatter which do involve a change in composition. _ ~Two factors in all changes. ~ In all the changes which matter can undergo, whether physical or chemical, two factors must be taken into account, namely, _energy_ and _matter_. ~Energy. ~ It is a familiar fact that certain bodies have the power to dowork. Thus water falling from a height upon a water wheel turns thewheel and in this way does the work of the mills. Magnetized ironattracts iron to itself and the motion of the iron as it moves towardsthe magnet can be made to do work. When coal is burned it causes theengine to move and transports the loaded cars from place to place. Whena body has this power to do work it is said to possess energy. ~Law of conservation of energy. ~ Careful experiments have shown that whenone body parts with its energy the energy is not destroyed but istransferred to another body or system of bodies. Just as energy cannotbe destroyed, neither can it be created. If one body gains a certainamount of energy, some other body has lost an equivalent amount. Thesefacts are summed up in the law of conservation of energy which may bestated thus: _While energy can be changed from one form into another, itcannot be created or destroyed. _ ~Transformations of energy. ~ Although energy can neither be created nordestroyed, it is evident that it may assume many different forms. Thusthe falling water may turn the electric generator and produce a currentof electricity. The energy lost by the falling water is thus transformedinto the energy of the electric current. This in turn may be changedinto the energy of motion, as when the current is used for propellingthe cars, or into the energy of heat and light, as when it is used forheating and lighting the cars. Again, the energy of coal may beconverted into energy of heat and subsequently of motion, as when it isused as a fuel in steam engines. Since the energy possessed by coal only becomes available when the coalis made to undergo a chemical change, it is sometimes called _chemicalenergy_. It is this form of energy in which we are especially interestedin the study of chemistry. ~Matter. ~ Matter may be defined as that which occupies space and possessesweight. Like energy, matter may be changed oftentimes from one form intoanother; and since in these transformations all the other physicalproperties of a substance save weight are likely to change, the inquiryarises, Does the weight also change? Much careful experimenting hasshown that it does not. The weight of the products formed in any changein matter always equals the weight of the substances undergoing change. ~Law of conservation of matter. ~ The important truth just stated isfrequently referred to as the law of conservation of matter, and thislaw may be briefly stated thus: _Matter can neither be created nordestroyed, though it can be changed from one form into another. _ ~Classification of matter. ~ At first sight there appears to be no limit tothe varieties of matter of which the world is made. For convenience instudy we may classify all these varieties under three heads, namely, _mechanical mixtures_, _chemical compounds_, and _elements_. [Illustration: Fig. 1] ~Mechanical mixtures. ~ If equal bulks of common salt and iron filings arethoroughly mixed together, a product is obtained which, judging by itsappearance, is a new substance. If it is examined more closely, however, it will be seen to be merely a mixture of the salt and iron, each ofwhich substances retains its own peculiar properties. The mixture tastesjust like salt; the iron particles can be seen and their grittycharacter detected. A magnet rubbed in the mixture draws out the ironjust as if the salt were not there. On the other hand, the salt can beseparated from the iron quite easily. Thus, if several grams of themixture are placed in a test tube, and the tube half filled with waterand thoroughly shaken, the salt dissolves in the water. The ironparticles can then be filtered from the liquid by pouring the entiremixture upon a piece of filter paper folded so as to fit into theinterior of a funnel (Fig. 1). The paper retains the solid but allowsthe clear liquid, known as the _filtrate_, to drain through. The ironparticles left upon the filter paper will be found to be identical withthe original iron. The salt can be recovered from the filtrate byevaporation of the water. To accomplish this the filtrate is poured intoa small evaporating dish and gently heated (Fig. 2) until the water hasdisappeared, or _evaporated_. The solid left in the dish is identical inevery way with the original salt. Both the iron and the salt have thusbeen recovered in their original condition. It is evident that no newsubstance has been formed by rubbing the salt and iron together. Theproduct is called a _mechanical mixture_. Such mixtures are very commonin nature, almost all minerals, sands, and soils being examples of thisclass of substances. It is at once apparent that there is no lawregulating the composition of a mechanical mixture, and no two mixturesare likely to have exactly the same composition. The ingredients of amechanical mixture can usually be separated by mechanical means, such assifting, sorting, magnetic attraction, or by dissolving one constituentand leaving the other unchanged. [Illustration: Fig. 2] DEFINITION: _A mechanical mixture is one in which the constituentsretain their original properties, no chemical action having taken placewhen they were brought together. _ ~Chemical compounds. ~ If iron filings and powdered sulphur are thoroughlyground together in a mortar, a yellowish-green substance results. Itmight easily be taken to be a new body; but as in the case of the ironand salt, the ingredients can readily be separated. A magnet draws outthe iron. Water does not dissolve the sulphur, but other liquids do, as, for example, the liquid called carbon disulphide. When the mixture istreated with carbon disulphide the iron is left unchanged, and thesulphur can be obtained again, after filtering off the iron, byevaporating the liquid. The substance is, therefore, a mechanicalmixture. If now a new portion of the mixture is placed in a dry test tube andcarefully heated in the flame of a Bunsen burner, as shown in Fig. 3, astriking change takes place. The mixture begins to glow at some point, the glow rapidly extending throughout the whole mass. If the test tubeis now broken and the product examined, it will be found to be a hard, black, brittle substance, in no way recalling the iron or the sulphur. The magnet no longer attracts it; carbon disulphide will not dissolvesulphur from it. It is a new substance with new properties, resultingfrom the chemical union of iron and sulphur, and is called ironsulphide. Such substances are called _chemical compounds_, and differfrom mechanical mixtures in that the substances producing them losetheir own characteristic properties. We shall see later that the twoalso differ in that the composition of a chemical compound never varies. [Illustration: Fig. 3] DEFINITION: _A chemical compound is a substance the constituents ofwhich have lost their own characteristic properties, and which cannot beseparated save by a chemical change. _ ~Elements. ~ It has been seen that iron sulphide is composed of twoentirely different substances, --iron and sulphur. The question arises, Do these substances in turn contain other substances, that is, are theyalso chemical compounds? Chemists have tried in a great many ways todecompose them, but all their efforts have failed. Substances which haveresisted all efforts to decompose them into other substances are called_elements_. It is not always easy to prove that a given substance isreally an element. Some way as yet untried may be successful indecomposing it into other simpler forms of matter, and the supposedelement will then prove to be a compound. Water, lime, and many otherfamiliar compounds were at one time thought to be elements. DEFINITION: _An element is a substance which cannot be separated intosimpler substances by any known means. _ ~Kinds of matter. ~ While matter has been grouped in three classes for thepurpose of study, it will be apparent that there are really but twodistinct kinds of matter, namely, compounds and elements. A mechanicalmixture is not a third distinct kind of matter, but is made up ofvarying quantities of either compounds or elements or both. ~Alchemy. ~ In olden times it was thought that some way could be found tochange one element into another, and a great many efforts were made toaccomplish this transformation. Most of these efforts were directedtoward changing the commoner metals into gold, and many fanciful waysfor doing this were described. The chemists of that time were called_alchemists_, and the art which they practiced was called _alchemy_. Thealchemists gradually became convinced that the only way common metalscould be changed into gold was by the wonderful power of a magicsubstance which they called the _philosopher's stone_, which wouldaccomplish this transformation by its mere touch and would in additiongive perpetual youth to its fortunate possessor. No one has ever foundsuch a stone, and no one has succeeded in changing one metal intoanother. ~Number of elements. ~ The number of substances now considered to beelements is not large--about eighty in all. Many of these are rare, andvery few of them make any large fraction of the materials in theearth's crust. Clarke gives the following estimate of the composition ofthe earth's crust: Oxygen 47. 0% Calcium 3. 5% Silicon 27. 9 Magnesium 2. 5 Aluminium 8. 1 Sodium 2. 7 Iron 4. 7 Potassium 2. 4 Other elements 1. 2% A complete list of the elements is given in the Appendix. In this listthe more common of the elements are marked with an asterisk. It is notnecessary to study more than a third of the total number of elements togain a very good knowledge of chemistry. ~Physical state of the elements. ~ About ten of the elements are gases atordinary temperatures. Two--mercury and bromine--are liquids. The othersare all solids, though their melting points vary through wide limits, from cęsium which melts at 26° to elements which do not melt save in theintense heat of the electric furnace. ~Occurrence of the elements. ~ Comparatively few of the elements occur asuncombined substances in nature, most of them being found in the form ofchemical compounds. When an element does occur by itself, as is the casewith gold, we say that it occurs in the _free state_ or _native_; whenit is combined with other substances in the form of compounds, we saythat it occurs in the _combined state_, or _in combination_. In thelatter case there is usually little about the compound to suggest thatthe element is present in it; for we have seen that elements lose theirown peculiar properties when they enter into combination with otherelements. It would never be suspected, for example, that the reddish, earthy-looking iron ore contains iron. ~Names of elements. ~ The names given to the elements have been selected ina great many different ways. (1) Some names are very old and theiroriginal meaning is obscure. Such names are iron, gold, and copper. (2)Many names indicate some striking physical property of the element. Thename bromine, for example, is derived from a Greek word meaning astench, referring to the extremely unpleasant odor of the substance. Thename iodine comes from a word meaning violet, alluding to the beautifulcolor of iodine vapor. (3) Some names indicate prominent chemicalproperties of the elements. Thus, nitrogen means the producer of niter, nitrogen being a constituent of niter or saltpeter. Hydrogen means waterformer, signifying its presence in water. Argon means lazy or inert, theelement being so named because of its inactivity. (4) Other elements arenamed from countries or localities, as germanium and scandium. ~Symbols. ~ In indicating the elements found in compounds it isinconvenient to use such long names, and hence chemists have adopted asystem of abbreviations. These abbreviations are known as _symbols_, each element having a distinctive symbol. (1) Sometimes the initialletter of the name will suffice to indicate the element. Thus I standsfor iodine, C for carbon. (2) Usually it is necessary to add some othercharacteristic letter to the symbol, since several names may begin withthe same letter. Thus C stands for carbon, Cl for chlorine, Cd forcadmium, Ce for cerium, Cb for columbium. (3) Sometimes the symbol is anabbreviation of the old Latin name. In this way Fe (ferrum) indicatesiron, Cu (cuprum), copper, Au (aurum), gold. The symbols are included inthe list of elements given in the Appendix. They will become familiarthrough constant use. ~Chemical affinity the cause of chemical combination. ~ The agency whichcauses substances to combine and which holds them together when combinedis called _chemical affinity_. The experiments described in thischapter, however, show that heat is often necessary to bring aboutchemical action. The distinction between the cause producing chemicalaction and the circumstances favoring it must be clearly made. Chemicalaffinity is always the cause of chemical union. Many agencies may makeit possible for chemical affinity to act by overcoming circumstanceswhich stand in its way. Among these agencies are heat, light, andelectricity. As a rule, solution also promotes action between twosubstances. Sometimes these agencies may overcome chemical attractionand so occasion the decomposition of a compound. EXERCISES 1. To what class of changes do the following belong? (a) The meltingof ice; (b) the souring of milk; (c) the burning of a candle; (d)the explosion of gunpowder; (e) the corrosion of metals. What testquestion must be applied in each of the above cases? 2. Give two additional examples (a) of chemical changes; (b) ofphysical changes. 3. Is a chemical change always accompanied by a physical change? Is aphysical change always accompanied by a chemical change? 4. Give two or more characteristics of a chemical change. 5. (a) When a given weight of water freezes, does it absorb or evolveheat? (b) When the resulting ice melts, is the total heat change thesame or different from that of freezing? 6. Give three examples of each of the following: (a) mechanicalmixtures; (b) chemical compounds; (c) elements. 7. Give the derivation of the names of the following elements: thorium, gallium, selenium, uranium. (Consult dictionary. ) 8. Give examples of chemical changes which are produced through theagency of heat; of light; of electricity. CHAPTER II OXYGEN ~History. ~ The discovery of oxygen is generally attributed to the Englishchemist Priestley, who in 1774 obtained the element by heating acompound of mercury and oxygen, known as red oxide of mercury. It isprobable, however, that the Swedish chemist Scheele had previouslyobtained it, although an account of his experiments was not publisheduntil 1777. The name oxygen signifies acid former. It was given to theelement by the French chemist Lavoisier, since he believed that allacids owe their characteristic properties to the presence of oxygen. This view we now know to be incorrect. ~Occurrence. ~ Oxygen is by far the most abundant of all the elements. Itoccurs both in the free and in the combined state. In the free state itoccurs in the air, 100 volumes of dry air containing about 21 volumes ofoxygen. In the combined state it forms eight ninths of water and nearlyone half of the rocks composing the earth's crust. It is also animportant constituent of the compounds which compose plant and animaltissues; for example, about 66% by weight of the human body is oxygen. ~Preparation. ~ Although oxygen occurs in the free state in the atmosphere, its separation from the nitrogen and other gases with which it is mixedis such a difficult matter that in the laboratory it has been found moreconvenient to prepare it from its compounds. The most important of thelaboratory methods are the following: 1. _Preparation from water. _ Water is a compound, consisting of 11. 18%hydrogen and 88. 82% oxygen. It is easily separated into theseconstituents by passing an electric current through it under suitableconditions. The process will be described in the chapter on water. Whilethis method of preparation is a simple one, it is not economical. 2. _Preparation from mercuric oxide. _ This method is of interest, sinceit is the one which led to the discovery of oxygen. The oxide, whichconsists of 7. 4% oxygen and 92. 6% mercury, is placed in a small, glasstest tube and heated. The compound is in this way decomposed intomercury which collects on the sides of the glass tube, forming a silverymirror, and oxygen which, being a gas, escapes from the tube. Thepresence of the oxygen is shown by lighting the end of a splint, extinguishing the flame and bringing the glowing coal into the mouth ofthe tube. The oxygen causes the glowing coal to burst into a flame. In a similar way oxygen may be obtained from its compounds with some of the other elements. Thus manganese dioxide, a black compound of manganese and oxygen, when heated to about 700°, loses one third of its oxygen, while barium dioxide, when heated, loses one half of its oxygen. 3. _Preparation from potassium chlorate (usual laboratory method). _Potassium chlorate is a white solid which consists of 31. 9% potassium, 28. 9% chlorine, and 39. 2% oxygen. When heated it undergoes a series ofchanges in which all the oxygen is finally set free, leaving a compoundof potassium and chlorine called potassium chloride. The change may berepresented as follows: /potassium\ | | (potassium / potassium \ (potassium { chlorine } = { } + oxygen | | chlorate) \ chlorine / chloride) \oxygen / [Illustration: JOSEPH PRIESTLEY (English) (1733-1804) School-teacher, theologian, philosopher, scientist; friend of BenjaminFranklin; discoverer of oxygen; defender of the phlogiston theory; thefirst to use mercury in a pneumatic trough, by which means he firstisolated in gaseous form hydrochloric acid, sulphur dioxide, andammonia] The evolution of the oxygen begins at about 400°. It has been found, however, that if the potassium chlorate is mixed with about one fourthits weight of manganese dioxide, the oxygen is given off at a much lowertemperature. Just how the manganese dioxide brings about this result isnot definitely known. The amount of oxygen obtained from a given weightof potassium chlorate is exactly the same whether the manganese dioxideis present or not. So far as can be detected the manganese dioxideundergoes no change. [Illustration: Fig. 4] ~Directions for preparing oxygen. ~ The manner of preparing oxygen frompotassium chlorate is illustrated in the accompanying diagram (Fig. 4). A mixture consisting of one part of manganese dioxide and four parts ofpotassium chlorate is placed in the flask A and gently heated. Theoxygen is evolved and escapes through the tube B. It is collected bybringing over the end of the tube the mouth of a bottle completelyfilled with water and inverted in a vessel of water, as shown in thefigure. The gas rises in the bottle and displaces the water. In thepreparation of large quantities of oxygen, a copper retort (Fig. 5) isoften substituted for the glass flask. [Illustration: Fig. 5] In the preparation of oxygen from potassium chlorate and manganesedioxide, the materials used must be pure, otherwise a violent explosionmay occur. The purity of the materials is tested by heating a smallamount of the mixture in a test tube. ~The collection of gases. ~ The method used for collecting oxygenillustrates the general method used for collecting such gases as areinsoluble in water or nearly so. The vessel C (Fig. 4), containing thewater in which the bottles are inverted, is called a _pneumatic trough. _ ~Commercial methods of preparation. ~ Oxygen can now be purchased storedunder great pressure in strong steel cylinders (Fig. 6). It is preparedeither by heating a mixture of potassium chlorate and manganese dioxide, or by separating it from the nitrogen and other gases with which it ismixed in the atmosphere. The methods employed for effecting thisseparation will be described in subsequent chapters. [Illustration: Fig. 6] ~Physical properties. ~ Oxygen is a colorless, odorless, tasteless gas, slightly heavier than air. One liter of it, measured at a temperature of0° and under a pressure of one atmosphere, weighs 1. 4285 g. , while undersimilar conditions one liter of air weighs 1. 2923 g. It is but slightlysoluble in water. Oxygen, like other gases, may be liquefied by applyingvery great pressure to the highly cooled gas. When the pressure isremoved the liquid oxygen passes again into the gaseous state, since itsboiling point under ordinary atmospheric pressure is -182. 5°. ~Chemical properties. ~ At ordinary temperatures oxygen is not very activechemically. Most substances are either not at all affected by it, or theaction is so slow as to escape notice. At higher temperatures, however, it is very active, and unites directly with most of the elements. Thisactivity may be shown by heating various substances until just ignitedand then bringing them into vessels of the gas, when they will burn withgreat brilliancy. Thus a glowing splint introduced into a jar of oxygenbursts into flame. Sulphur burns in the air with a very weak flame andfeeble light; in oxygen, however, the flame is increased in size andbrightness. Substances which readily burn in air, such as phosphorus, burn in oxygen with dazzling brilliancy. Even substances which burn inair with great difficulty, such as iron, readily burn in oxygen. The burning of a substance in oxygen is due to the rapid combination ofthe substance or of the elements composing it with the oxygen. Thus, when sulphur burns both the oxygen and sulphur disappear as such andthere is formed a compound of the two, which is an invisible gas, havingthe characteristic odor of burning sulphur. Similarly, phosphorus onburning forms a white solid compound of phosphorus and oxygen, whileiron forms a reddish-black compound of iron and oxygen. ~Oxidation. ~ The term _oxidation_ is applied to the chemical change whichtakes place when a substance, or one of its constituent parts, combineswith oxygen. This process may take place rapidly, as in the burning ofphosphorus, or slowly, as in the oxidation (or rusting) of iron whenexposed to the air. It is always accompanied by the liberation of heat. The amount of heat liberated by the oxidation of a definite weight ofany given substance is always the same, being entirely independent ofthe rapidity of the process. If the oxidation takes place slowly, theheat is generated so slowly that it is difficult to detect it. If theoxidation takes place rapidly, however, the heat is generated in such ashort interval of time that the substance may become white hot or burstinto a flame. ~Combustion; kindling temperature. ~ When oxidation takes place so rapidlythat the heat generated is sufficient to cause the substance to glow orburst into a flame the process is called _combustion_. In order that anysubstance may undergo combustion, it is necessary that it should beheated to a certain temperature, known as the _kindling temperature. _This temperature varies widely for different bodies, but is alwaysdefinite for the same body. Thus the kindling temperature of phosphorusis far lower than that of iron, but is definite for each. When anyportion of a substance is heated until it begins to burn the combustionwill continue without the further application of heat, provided the heatgenerated by the process is sufficient to bring other parts of thesubstance to the kindling temperature. On the other hand, if the heatgenerated is not sufficient to maintain the kindling temperature, combustion ceases. ~Oxides. ~ The compounds formed by the oxidation of any element are called_oxides_. Thus in the combustion of sulphur, phosphorus, and iron, thecompounds formed are called respectively oxide of sulphur, oxide ofphosphorus, and oxide of iron. In general, then, _an oxide is a compoundof oxygen with another element_. A great many substances of this classare known; in fact, the oxides of all the common elements have beenprepared, with the exception of those of fluorine and bromine. Some ofthese are familiar compounds. Water, for example, is an oxide ofhydrogen, and lime an oxide of the metal calcium. ~Products of combustion. ~ The particular oxides formed by the combustionof any substance are called _products of combustion_ of that substance. Thus oxide of sulphur is the product of the combustion of sulphur; oxideof iron is the product of the combustion of iron. It is evident that theproducts of the combustion of any substance must weigh more than theoriginal substance, the increase in weight corresponding to the amountof oxygen taken up in the act of combustion. For example, when ironburns the oxide of iron formed weighs more than the original iron. In some cases the products of combustion are invisible gases, so thatthe substance undergoing combustion is apparently destroyed. Thus, whena candle burns it is consumed, and so far as the eye can judge nothingis formed during combustion. That invisible gases are formed, however, and that the weight of these is greater than the weight of the candlemay be shown by the following experiment. [Illustration: Fig. 7] A lamp chimney is filled with sticks of the compound known as sodium hydroxide (caustic soda), and suspended from the beam of the balance, as shown in Fig. 7. A piece of candle is placed on the balance pan so that the wick comes just below the chimney, and the balance is brought to a level by adding weights to the other pan. The candle is then lighted. The products formed pass up through the chimney and are absorbed by the sodium hydroxide. Although the candle burns away, the pan upon which it rests slowly sinks, showing that the combustion is attended by an increase in weight. ~Combustion in air and in oxygen. ~ Combustion in air and in oxygen differs only in rapidity, the products formed being exactly the same. That the process should take place less rapidly in the former is readily understood, for the air is only about one fifth oxygen, the remaining four fifths being inert gases. Not only is less oxygen available, but much of the heat is absorbed in raising the temperature of the inert gases surrounding the substance undergoing combustion, and the temperature reached in the combustion is therefore less. ~Phlogiston theory of combustion. ~ The French chemist Lavoisier (1743-1794), who gave to oxygen its name was the first to show that combustion is due to union with oxygen. Previous to his time combustion was supposed to be due to the presence of a substance or principle called _phlogiston_. One substance was thought to be more combustible than another because it contained more phlogiston. Coal, for example, was thought to be very rich in phlogiston. The ashes left after combustion would not burn because all the phlogiston had escaped. If the phlogiston could be restored in any way, the substance would then become combustible again. Although this view seems absurd to us in the light of our present knowledge, it formerly had general acceptance. The discovery of oxygen led Lavoisier to investigate the subject, and through his experiments he arrived at the true explanation of combustion. The discovery of oxygen together with the part it plays in combustion is generally regarded as the most important discovery in the history of chemistry. It marked the dawn of a new period in the growth of the science. ~Combustion in the broad sense. ~ According to the definition given above, the presence of oxygen is necessary for combustion. The term issometimes used, however, in a broader sense to designate any chemicalchange attended by the evolution of heat and light. Thus iron andsulphur, or hydrogen and chlorine under certain conditions, will combineso rapidly that light is evolved, and the action is called a combustion. Whenever combustion takes place in the air, however, the process is oneof oxidation. ~Spontaneous combustion. ~ The temperature reached in a given chemical action, such as oxidation, depends upon the rate at which the reaction takes place. This rate is usually increased by raising the temperature of the substances taking part in the action. When a slow oxidation takes place under such conditions that the heat generated is not lost by being conducted away, the temperature of the substance undergoing oxidation is raised, and this in turn hastens the rate of oxidation. The rise in temperature may continue in this way until the kindling temperature of the substance is reached, when combustion begins. Combustion occurring in this way is called _spontaneous combustion_. Certain oils, such as the linseed oil used in paints, slowly undergo oxidation at ordinary temperatures, and not infrequently the origin of fires has been traced to the spontaneous combustion of oily rags. The spontaneous combustion of hay has been known to set barns on fire. Heaps of coal have been found to be on fire when spontaneous combustion offered the only possible explanation. ~Importance of oxygen. ~ 1. Oxygen is essential to life. Among livingorganisms only certain minute forms of plant life can exist without it. In the process of respiration the air is taken into the lungs where acertain amount of oxygen is absorbed by the blood. It is then carried toall parts of the body, oxidizing the worn-out tissues and changing theminto substances which may readily be eliminated from the body. The heatgenerated by this oxidation is the source of the heat of the body. Thesmall amount of oxygen which water dissolves from the air supports allthe varied forms of aquatic animals. 2. Oxygen is also essential to decay. The process of decay is really akind of oxidation, but it will only take place in the presence ofcertain minute forms of life known as bacteria. Just how these assist inthe oxidation is not known. By this process the dead products of animaland vegetable life which collect on the surface of the earth are slowlyoxidized and so converted into harmless substances. In this way oxygenacts as a great purifying agent. 3. Oxygen is also used in the treatment of certain diseases in which thepatient is unable to inhale sufficient air to supply the necessaryamount of oxygen. OZONE ~Preparation. ~ When electric sparks are passed through oxygen or air asmall percentage of the oxygen is converted into a substance called_ozone_, which differs greatly from oxygen in its properties. The samechange can also be brought about by certain chemical processes. Thus, ifsome pieces of phosphorus are placed in a bottle and partially coveredwith water, the presence of ozone may soon be detected in the aircontained in the bottle. The conversion of oxygen into ozone is attendedby a change in volume, 3 volumes of oxygen forming 2 volumes of ozone. If the resulting ozone is heated to about 300°, the reverse changetakes place, the 2 volumes of ozone being changed back into 3 volumes ofoxygen. It is possible that traces of ozone exist in the atmosphere, although its presence there has not been definitely proved, the testsformerly used for its detection having been shown to be unreliable. ~Properties. ~ As commonly prepared, ozone is mixed with a large excess ofoxygen. It is possible, however, to separate the ozone and thus obtainit in pure form. The gas so obtained has the characteristic odor noticedabout electrical machines when in operation. By subjecting it to greatpressure and a low temperature, the gas condenses to a bluish liquid, boiling at -119°. When unmixed with other gases ozone is very explosive, changing back into oxygen with the liberation of heat. Its chemicalproperties are similar to those of oxygen except that it is far moreactive. Air or oxygen containing a small amount of ozone is now used inplace of oxygen in certain manufacturing processes. ~The difference between oxygen and ozone. ~ Experiments show that inchanging oxygen into ozone no other kind of matter is either added tothe oxygen or withdrawn from it. The question arises then, How can weaccount for the difference in their properties? It must be rememberedthat in all changes we have to take into account _energy_ as well as_matter_. By changing the amount of energy in a substance we change itsproperties. That oxygen and ozone contain different amounts of energymay be shown in a number of ways; for example, by the fact that theconversion of ozone into oxygen is attended by the liberation of heat. The passage of the electric sparks through oxygen has in some waychanged the energy content of the element and thus it has acquired newproperties. _Oxygen and ozone must, therefore, be regarded as identicalso far as the kind of matter of which they are composed is concerned. Their different properties are due to their different energy contents. _ ~Allotropic states or forms of matter. ~ Other elements besides oxygen mayexist in more than one form. These different forms of the same elementare called _allotropic states_ or _forms_ of the element. These formsdiffer not only in physical properties but also in their energycontents. Elements often exist in a variety of forms which look quitedifferent. These differences may be due to accidental causes, such asthe size or shape of the particles or the way in which the element wasprepared. Only such forms, however, as have different energy contentsare properly called allotropic forms. MEASUREMENT OF GAS VOLUMES ~Standard conditions. ~ It is a well-known fact that the volume occupied bya definite weight of any gas can be altered by changing the temperatureof the gas or the pressure to which it is subjected. In measuring thevolume of gases it is therefore necessary, for the sake of accuracy, toadopt some standard conditions of temperature and pressure. Theconditions agreed upon are (1) a temperature of 0°, and (2) a pressureequal to the average pressure exerted by the atmosphere at the sealevel, that is, 1033. 3 g. Per square centimeter. These conditions oftemperature and pressure are known as the _standard conditions_, andwhen the volume of a gas is given it is understood that the measurementwas made under these conditions, unless it is expressly statedotherwise. For example, the weight of a liter of oxygen has been givenas 1. 4285 g. This means that one liter of oxygen, measured at atemperature of 0° and under a pressure of 1033. 3 g. Per squarecentimeter, weighs 1. 4285 g. The conditions which prevail in the laboratory are never the standardconditions. It becomes necessary, therefore, to find a way to calculatethe volume which a gas will occupy under standard conditions from thevolume which it occupies under any other conditions. This may be done inaccordance with the following laws. ~Law of Charles. ~ This law expresses the effect which a change in thetemperature of a gas has upon its volume. It may be stated as follows:_For every degree the temperature of a gas rises above zero the volumeof the gas is increased by 1/273 of the volume which it occupies atzero; likewise for every degree the temperature of the gas falls belowzero the volume of the gas is decreased by 1/273 of the volume which itoccupies at zero, provided in both cases that the pressure to which thegas is subjected remains constant. _ If V represents the volume of gas at 0°, then the volume at 1° will beV + 1/273 V; at 2° it will be V + 2/273 V; or, in general, thevolume v, at the temperature t, will be expressed by the formula (1) v = V + t/273 V, or (2) v = V(1 + (t/273)). Since 1/273 = 0. 00366, the formula may be written (3) v = V(1 + 0. 00366t). Since the value of V (volume under standard conditions) is the oneusually sought, it is convenient to transpose the equation to thefollowing form: (4) V = v/(1 + 0. 00366t). The following problem will serve as an illustration of the applicationof this equation. The volume of a gas at 20° is 750 cc. ; find the volume it will occupy at0°, the pressure remaining constant. In this case, v = 750 cc. And t = 20. By substituting these values, equation (4) becomes V = 750/(1 + 0. 00366 × 20) = 698. 9 cc. ~Law of Boyle. ~ This law expresses the relation between the volumeoccupied by a gas and the pressure to which it is subjected. It may bestated as follows: _The volume of a gas is inversely proportional to thepressure under which it is measured, provided the temperature of the gasremains constant. _ If V represents the volume when subjected to a pressure P and vrepresents its volume when the pressure is changed to p, then, inaccordance with the above law, V : v :: p : P, or VP = vp. In other words, for a given weight of a gas the product of the numbersrepresenting its volume and the pressure to which it is subjected is aconstant. Since the pressure of the atmosphere at any point is indicated by thebarometric reading, it is convenient in the solution of the problems tosubstitute the latter for the pressure measured in grams per squarecentimeter. The average reading of the barometer at the sea level is 760mm. , which corresponds to a pressure of 1033. 3 g. Per square centimeter. The following problem will serve as an illustration of the applicationof Boyle's law. A gas occupies a volume of 500 cc. In a laboratory where the barometricreading is 740 mm. What volume would it occupy if the atmosphericpressure changed so that the reading became 750 mm. ? Substituting the values in the equation VP = vp, we have 500 × 740 =v × 750, or v = 493. 3 cc. ~Variations in the volume of a gas due to changes both in temperature andpressure. ~ Inasmuch as corrections must be made as a rule for bothtemperature and pressure, it is convenient to combine the equationsgiven above for the corrections for each, so that the two correctionsmay be made in one operation. The following equation is thus obtained: (5) V_{s} = vp/(760(1 + 0. 00366t)), in which V_{s} represents the volume of a gas under standardconditions and v, p, and t the volume, pressure, and temperaturerespectively at which the gas was actually measured. The following problem will serve to illustrate the application of thisequation. A gas having a temperature of 20° occupies a volume of 500 cc. Whensubjected to a pressure indicated by a barometric reading of 740 mm. What volume would this gas occupy under standard conditions? In this problem v = 500, p = 740, and t = 20. Substituting thesevalues in the above equation, we get V_{s} = (500 × 740)/(760 (1 + 0. 00366 × 20)) = 453. 6 cc. [Illustration: Fig. 8] ~Variations in the volume of a gas due to the pressure of aqueous vapor. ~In many cases gases are collected over water, as explained under thepreparation of oxygen. In such cases there is present in the gas acertain amount of water vapor. This vapor exerts a definite pressure, which acts in opposition to the atmospheric pressure and which thereforemust be subtracted from the latter in determining the effective pressureupon the gas. Thus, suppose we wish to determine the pressure to whichthe gas in tube A (Fig. 8) is subjected. The tube is raised or lowereduntil the level of the water inside and outside the tube is the same. The atmosphere presses down upon the surface of the water (as indicatedby the arrows), thus forcing the water upward within the tube with apressure equal to the atmospheric pressure. The full force of thisupward pressure, however, is not spent in compressing the gas within thetube, for since it is collected over water it contains a certain amountof water vapor. This water vapor exerts a pressure (as indicated by thearrow within the tube) in opposition to the upward pressure. It isplain, therefore, that the effective pressure upon the gas is equal tothe atmospheric pressure less the pressure exerted by the aqueous vapor. The pressure exerted by the aqueous vapor increases with thetemperature. The figures representing the extent of this pressure (oftencalled the _tension of aqueous vapor_) are given in the Appendix. Theyexpress the pressure or tension in millimeters of mercury, just as theatmospheric pressure is expressed in millimeters of mercury. Representing the pressure of the aqueous vapor by a, formula (5)becomes (6) V_{s} = v(p - a)/(760(1 + 0. 00366t)). The following problem will serve to illustrate the method of applyingthe correction for the pressure of the aqueous vapor. The volume of a gas measured over water in a laboratory where thetemperature is 20° and the barometric reading is 740 mm. Is 500 cc. Whatvolume would this occupy under standard conditions? The pressure exerted by the aqueous vapor at 20° (see table in Appendix)is equal to the pressure exerted by a column of mercury 17. 4 mm. Inheight. Substituting the values of v, t, p, and a in formula(6), we have (6) V_{s} = 500(740 - 17. 4)/(760(1 + 0. 00366 × 20)) = 442. 9 cc. ~Adjustment of tubes before reading gas volumes. ~ In measuring the volumesof gases collected in graduated tubes or other receivers, over a liquidas illustrated in Fig. 8, the reading should be taken after raising orlowering the tube containing the gas until the level of the liquidinside and outside the tube is the same; for it is only under theseconditions that the upward pressure within the tube is the same as theatmospheric pressure. EXERCISES 1. What is the meaning of the following words? phlogiston, ozone, phosphorus. (Consult dictionary. ) 2. Can combustion take place without the emission of light? 3. Is the evolution of light always produced by combustion? 4. (a) What weight of oxygen can be obtained from 100 g. Of water?(b) What volume would this occupy under standard conditions? 5. (a) What weight of oxygen can be obtained from 500g. Of mercuricoxide? (b) What volume would this occupy under standard conditions? 6. What weight of each of the following compounds is necessary toprepare 50 l. Of oxygen? (a) water; (b) mercuric oxide; (c)potassium chlorate. 7. Reduce the following volumes to 0°, the pressure remaining constant:(a) 150 cc. At 10°; (b) 840 cc. At 273°. 8. A certain volume of gas is measured when the temperature is 20°. Atwhat temperature will its volume be doubled? 9. Reduce the following volumes to standard conditions of pressure, thetemperature remaining constant: (a) 200 cc. At 740 mm. ; (b) 500 l. At 380 mm. 10. What is the weight of 1 l. Of oxygen when the pressure is 750 mm. And the temperature 0°? 11. Reduce the following volumes to standard conditions of temperatureand pressure: (a) 340 cc. At 12° and 753 mm; (b) 500 cc. At 15° and740 mm. 12. What weight of potassium chlorate is necessary to prepare 250 l. Ofoxygen at 20° and 750 mm. ? 13. Assuming the cost of potassium chlorate and mercuric oxide to berespectively $0. 50 and $1. 50 per kilogram, calculate the cost ofmaterials necessary for the preparation of 50 l. Of oxygen from each ofthe above compounds. 14. 100 g. Of potassium chlorate and 25 g. Of manganese dioxide wereheated in the preparation of oxygen. What products were left in theflask, and how much of each was present? CHAPTER III HYDROGEN ~Historical. ~ The element hydrogen was first clearly recognized as adistinct substance by the English investigator Cavendish, who in 1766obtained it in a pure state, and showed it to be different from theother inflammable airs or gases which had long been known. Lavoisiergave it the name hydrogen, signifying water former, since it had beenfound to be a constituent of water. ~Occurrence. ~ In the free state hydrogen is found in the atmosphere, butonly in traces. In the combined state it is widely distributed, being aconstituent of water as well as of all living organisms, and theproducts derived from them, such as starch and sugar. About 10% of thehuman body is hydrogen. Combined with carbon, it forms the substanceswhich constitute petroleum and natural gas. It is an interesting fact that while hydrogen in the free state occurs only in traces on the earth, it occurs in enormous quantities in the gaseous matter surrounding the sun and certain other stars. ~Preparation from water. ~ Hydrogen can be prepared from water by severalmethods, the most important of which are the following. 1. _By the electric current. _ As has been indicated in the preparationof oxygen, water is easily separated into its constituents, hydrogen andoxygen, by passing an electric current through it under certainconditions. 2. _By the action of certain metals. _ When brought into contact withcertain metals under appropriate conditions, water gives up a portionor the whole of its hydrogen, its place being taken by the metal. In thecase of a few of the metals this change occurs at ordinary temperatures. Thus, if a bit of sodium is thrown on water, an action is seen to takeplace at once, sufficient heat being generated to melt the sodium, whichruns about on the surface of the water. The change which takes placeconsists in the displacement of one half of the hydrogen of the water bythe sodium, and may be represented as follows: _ _ _ _ | hydrogen | | sodium |sodium + | hydrogen |(water) = | hydrogen |(sodium hydroxide) + hydrogen |_oxygen _| |_oxygen _| The sodium hydroxide formed is a white solid which remains dissolved inthe undecomposed water, and may be obtained by evaporating the solutionto dryness. The hydrogen is evolved as a gas and may be collected bysuitable apparatus. Other metals, such as magnesium and iron, decompose water rapidly, butonly at higher temperatures. When steam is passed over hot iron, forexample, the iron combines with the oxygen of the steam, thus displacingthe hydrogen. Experiments show that the change may be represented asfollows: _ _ | hydrogen | _ _ _ _iron + | hydrogen |(water) = | iron |(iron oxide) + | hydrogen | |_oxygen _| |_oxygen _| |_hydrogen_| The iron oxide formed is a reddish-black compound, identical with thatobtained by the combustion of iron in oxygen. ~Directions for preparing hydrogen by the action of steam on iron. ~ The apparatus used in the preparation of hydrogen from iron and steam is shown in Fig. 9. A porcelain or iron tube B, about 50 cm. In length and 2 cm. Or 3 cm. In diameter, is partially filled with fine iron wire or tacks and connected as shown in the figure. The tube B is heated, slowly at first, until the iron is red-hot. Steam is then conducted through the tube by boiling the water in the flask A. The hot iron combines with the oxygen in the steam, setting free the hydrogen, which is collected over water. The gas which first passes over is mixed with the air previously contained in the flask and tube, and is allowed to escape, _since a mixture of hydrogen with oxygen or air explodes violently when brought in contact with a flame_. It is evident that the flask A must be disconnected from the tube before the heat is withdrawn. That the gas obtained is different from air and oxygen may be shown by holding a bottle of it mouth downward and bringing a lighted splint into it. The hydrogen is ignited and burns with an almost colorless flame. [Illustration Fig. 9] ~Preparation from acids~ (_usual laboratory method_). While hydrogen canbe prepared from water, either by the action of the electric current orby the action of certain metals, these methods are not economical andare therefore but little used. In the laboratory hydrogen is generallyprepared from compounds known as acids, all of which contain hydrogen. When acids are brought in contact with certain metals, the metalsdissolve and set free the hydrogen of the acid. Although this reactionis a quite general one, it has been found most convenient in preparinghydrogen by this method to use either zinc or iron as the metal andeither hydrochloric or sulphuric acid as the acid. Hydrochloric acid isa compound consisting of 2. 77% hydrogen and 97. 23% chlorine, whilesulphuric acid consists of 2. 05% hydrogen, 32. 70% sulphur, and 65. 25%oxygen. The changes which take place in the preparation of hydrogen from zincand sulphuric acid (diluted with water) may be represented as follows: _ _ _ _ | hydrogen |(sulphuric | zinc |(zinczinc + | sulphur | acid) = | sulphur | sulphate) + hydrogen |_oxygen _| |_oxygen _| In other words, the zinc has taken the place of the hydrogen insulphuric acid. The resulting compound contains zinc, sulphur, andoxygen, and is known as zinc sulphate. This remains dissolved in thewater present in the acid. It may be obtained in the form of a whitesolid by evaporating the liquid left after the metal has passed intosolution. When zinc and hydrochloric acid are used the following changes takeplace: _ _ _ _ | hydrogen |(hydrochloric | zinc |(zinczinc + |_chlorine_| acid) = |_chlorine_| chloride) + hydrogen When iron is used the changes which take place are exactly similar tothose just given for zinc. [Illustration Fig. 10. ] ~Directions for preparing hydrogen from acids. ~ The preparation of hydrogen from acids is carried out in the laboratory as follows: The metal is placed in a flask or wide-mouthed bottle A (Fig. 10) and the acid is added slowly through the funnel tube B. The metal dissolves in the acid, while the hydrogen which is liberated escapes through the exit tube C and is collected over water. It is evident that the hydrogen which passes over first is mixed with the air from the bottle A. Hence care must be taken not to bring a flame near the exit tube, since, as has been stated previously, such a mixture explodes with great violence when brought in contact with a flame. ~Precautions. ~ Both sulphuric acid and zinc, if impure, are likely to contain small amounts of arsenic. Such materials should not be used in preparing hydrogen, since the arsenic present combines with a portion of the hydrogen to form a very poisonous gas known as arsine. On the other hand, chemically pure sulphuric acid, i. E. Sulphuric acid that is entirely free from impurities, will not act upon chemically pure zinc. The reaction may be started, however, by the addition of a few drops of a solution of copper sulphate or platinum tetrachloride. ~Physical properties. ~ Hydrogen is similar to oxygen in that it is acolorless, tasteless, odorless gas. It is characterized by its extremelightness, being the lightest of all known substances. One liter of thegas weighs only 0. 08984 g. On comparing this weight with that of anequal volume of oxygen, viz. , 1. 4285 g. , the latter is found to be 15. 88times as heavy as hydrogen. Similarly, air is found to be 14. 38 times asheavy as hydrogen. Soap bubbles blown with hydrogen rapidly rise in theair. On account of its lightness it is possible to pour it upward fromone bottle into another. Thus, if the bottle A (Fig. 11) is filledwith hydrogen, placed mouth downward by the side of bottle _B_, filledwith air, and is then gradually inverted under B as indicated in thefigure, the hydrogen will flow upward into bottle _B_, displacing theair. Its presence in bottle B may then be shown by bringing a lightedsplint to the mouth of the bottle, when the hydrogen will be ignited bythe flame. It is evident, from this experiment, that in order to retainthe gas in an open bottle the bottle must be placed mouth downward. [Illustration Fig. 11] Hydrogen is far more difficult to liquefy than any other gas, with theexception of helium, a rare element recently found to exist in theatmosphere. The English scientist Dewar, however, in 1898 succeeded notonly in obtaining hydrogen in liquid state but also as a solid. Liquidhydrogen is colorless and has a density of only 0. 07. Its boiling pointunder atmospheric pressure is -252°. Under diminished pressure thetemperature has been reduced to -262°. The solubility of hydrogen inwater is very slight, being still less than that of oxygen. Pure hydrogen produces no injurious results when inhaled. Of course onecould not live in an atmosphere of the gas, since oxygen is essential torespiration. ~Chemical properties. ~ At ordinary temperatures hydrogen is not an activeelement. A mixture of hydrogen and chlorine, however, will combine withexplosive violence at ordinary temperature if exposed to the sunlight. The union can be brought about also by heating. The product formed ineither case is hydrochloric acid. Under suitable conditions hydrogencombines with nitrogen to form ammonia, and with sulphur to form thefoul-smelling gas, hydrogen sulphide. The affinity of hydrogen foroxygen is so great that a mixture of hydrogen and oxygen or hydrogenand air explodes with great violence when heated to the kindlingtemperature (about 612°). Nevertheless under proper conditions hydrogenmay be made to burn quietly in either oxygen or air. The resultinghydrogen flame is almost colorless and is very hot. The combustion ofthe hydrogen is, of course, due to its union with oxygen. The product ofthe combustion is therefore a compound of hydrogen and oxygen. That thiscompound is water may be shown easily by experiment. [Illustration Fig. 12] ~Directions for burning hydrogen in air. ~ The combustion of hydrogen in air may be carried out safely as follows: The hydrogen is generated in the bottle A (Fig. 12), is dried by conducting it through the tube X, filled with some substance (generally calcium chloride) which has a great attraction for moisture, and escapes through the tube T, the end of which is drawn out to a jet. The hydrogen first liberated mixes with the air contained in the generator. If a flame is brought near the jet before this mixture has all escaped, a violent and very dangerous explosion results, since the entire apparatus is filled with the explosive mixture. On the other hand, if the flame is not applied until all the air has been expelled, the hydrogen is ignited and burns quietly, since only the small amount of it which escapes from the jet can come in contact with the oxygen of the air at any one time. By holding a cold, dry bell jar or bottle over the flame, in the manner shown in the figure, the steam formed by the combustion of the hydrogen is condensed, the water collecting in drops on the sides of the jar. ~Precautions. ~ In order to avoid danger it is absolutely necessary toprove that the hydrogen is free from air before igniting it. This can bedone by testing small amounts of the escaping gas. A convenient and safemethod of doing this is to fill a test tube with the gas by inverting itover the jet. The hydrogen, on account of its lightness, collects in thetube, displacing the air. After holding it over the jet for a fewmoments in order that it may be filled with the gas, the tube is gentlybrought, mouth downward, to the flame of a burner placed not nearer thanan arm's length from the jet. If the hydrogen is mixed with air a slightexplosion occurs, but if pure it burns quietly in the tube. Theoperation is repeated until the gas burns quietly, when the tube isquickly brought back over the jet for an instant, whereby the escapinghydrogen is ignited by the flame in the tube. [Illustration. Fig. 13] ~A mixture of hydrogen and oxygen is explosive. ~ That a mixture ofhydrogen and air is explosive may be shown safely as follows: A corkthrough which passes a short glass tube about 1 cm. In diameter isfitted air-tight into the tubule of a bell jar of 2 l. Or 3 l. Capacity. (A thick glass bottle with bottom removed may be used. ) The tube isclosed with a small rubber stopper and the bell jar filled withhydrogen, the gas being collected over water. When entirely filled withthe gas the jar is removed from the water and supported by blocks ofwood in order to leave the bottom of the jar open, as shown in Fig. 13. The stopper is now removed from the tube in the cork, and the hydrogen, which on account of its lightness escapes from the tube, is at oncelighted. As the hydrogen escapes, the air flows in at the bottom of thejar and mixes with the remaining portion of the hydrogen, so that amixture of the two soon forms, and a loud explosion results. Theexplosion is not dangerous, since the bottom of the jar is open, thusleaving room for the expansion of the hot gas. Since air is only one fifth oxygen, the remainder being inert gases, itmay readily be inferred that a mixture of hydrogen with pure oxygenwould be far more explosive than a mixture of hydrogen with air. Suchmixtures should not be made except in small quantities and byexperienced workers. ~Hydrogen does not support combustion. ~ While hydrogen is readilycombustible, it is not a supporter of combustion. In other words, substances will not burn in it. This may be shown by bringing a lightedcandle supported by a stiff wire into a bottle or cylinder of the puregas, as shown in Fig. 14. The hydrogen is ignited by the flame of thecandle and burns at the mouth of the bottle, where it comes in contactwith the oxygen in the air. When the candle is thrust up into the gas, its flame is extinguished on account of the absence of oxygen. If slowlywithdrawn, the candle is relighted as it passes through the layer ofburning hydrogen. [Illustration: Fig. 14] [Illustration: Fig. 15] ~Reduction. ~ On account of its great affinity for oxygen, hydrogen has thepower of abstracting it from many of its compounds. Thus, if a stream ofhydrogen, dried by passing through the tube B (Fig. 15), filled withcalcium chloride, is conducted through the tube C containing somecopper oxide, heated to a moderate temperature, the hydrogen abstractsthe oxygen from the copper oxide. The change may be represented asfollows: hydrogen + {copper} {hydrogen} {oxygen}(copper oxide) = {oxygen }(water) + copper The water formed collects in the cold portions of the tube C near itsend. In this experiment the copper oxide is said to undergo reduction. _Reduction may therefore be defined as the process of withdrawing oxygenfrom a compound. _ ~Relation of reduction to oxidation. ~ At the same time that the copperoxide is reduced it is clear that the hydrogen is oxidized, for itcombines with the oxygen given up by the copper oxide. The two processesare therefore very closely related, and it usually happens that when onesubstance is oxidized some other substance is reduced. That substancewhich gives up its oxygen is called an _oxidizing agent_, while thesubstance which unites with the oxygen is called a _reducing agent_. ~The oxyhydrogen blowpipe. ~ This is a form of apparatus used for burninghydrogen in pure oxygen. As has been previously stated, the flameproduced by the combustion of hydrogen in the air is very hot. It isevident that if pure oxygen is substituted for air, the temperaturereached will be much higher, since there are no inert gases to absorbthe heat. The oxyhydrogen blowpipe, used to effect this combination, consists of a small tube placed within a larger one, as shown in Fig. 16. [Illustration: Fig. 16] The hydrogen, stored under pressure, generally in steel cylinders, isfirst passed through the outer tube and ignited at the open end of thetube. The oxygen from a similar cylinder is then conducted through theinner tube, and mixes with the hydrogen at the end of the tube. In orderto produce the maximum heat, the hydrogen and oxygen must be admitted tothe blowpipe in the exact proportion in which they combine, viz. , 2volumes of hydrogen to 1 of oxygen, or by weight, 1 part of hydrogen to7. 94 parts of oxygen. The intensity of the heat may be shown by bringinginto the flame pieces of metal such as iron wire or zinc. These burnwith great brilliancy. Even platinum, having a melting point of 1779°, may be melted by the heat of the flame. While the oxyhydrogen flame is intensely hot, it is almost non-luminous. If directed against some infusible substance like ordinary lime (calciumoxide), the heat is so intense that the lime becomes incandescent andglows with a brilliant light. This is sometimes used as a source oflight, under the name of _Drummond_ or _lime light_. [Illustration: Fig. 17] ~The blast lamp. ~ A similar form of apparatus is commonly used in thelaboratory as a source of heat under the name _blast lamp_ (Fig. 17). This differs from the oxyhydrogen blowpipe only in the size of thetubes. In place of the hydrogen and oxygen the more accessible coal gasand air are respectively used. The former is composed largely of amixture of free hydrogen and gaseous compounds of carbon and hydrogen. While the temperature of the flame is not so high as that of theoxyhydrogen blowpipe, it nevertheless suffices for most chemicaloperations carried out in the laboratory. ~Uses of hydrogen. ~ On account of its cost, hydrogen is but little usedfor commercial purposes. It is sometimes used as a material for theinflation of balloons, but usually the much cheaper coal gas issubstituted for it. Even hot air is often used when the duration ofascension is very short. It has been used also as a source of heat andlight in the oxyhydrogen blowpipe. Where the electric current isavailable, however, this form of apparatus has been displaced almostentirely by the electric light and electric furnace, which are much moreeconomical and more powerful sources of light and heat. EXERCISES 1. Will a definite weight of iron decompose an unlimited weight ofsteam? 2. Why is oxygen passed through the inner tube of the oxyhydrogenblowpipe rather than the outer? 3. In Fig. 14, will the flame remain at the mouth of the tube? 4. From Fig. 15, suggest a way for determining experimentally thequantity of water formed in the reaction. 5. Distinguish clearly between the following terms: oxidation, reduction, combustion, and kindling temperature. 6. Is oxidation always accompanied by reduction? 7. What is the source of heat in the lime light? What is the exact useof lime in this instrument? 8. In Fig. 12, why is it necessary to dry the hydrogen by means of thecalcium chloride in the tube X? 9. At what pressure would the weight of 1 l. Of hydrogen be equal tothat of oxygen under standard conditions? 10. (a) What weight of hydrogen can be obtained from 150 g. Ofsulphuric acid? (b) What volume would this occupy under standardconditions? (c) The density of sulphuric acid is 1. 84. What volumewould the 150 g. Of the acid occupy? 11. How many liters of hydrogen can be obtained from 50 cc. Of sulphuricacid having a density of 1. 84? 12. Suppose you wish to fill five liter bottles with hydrogen, the gasto be collected over water in your laboratory, how many cubiccentimeters of sulphuric acid would be required? CHAPTER IV COMPOUNDS OF HYDROGEN AND OXYGEN; WATER AND HYDROGEN DIOXIDE WATER ~Historical. ~ Water was long regarded as an element. In 1781 Cavendishshowed that it is formed by the union of hydrogen and oxygen. Being abeliever in the phlogiston theory, however, he failed to interpret hisresults correctly. A few years later Lavoisier repeated Cavendish'sexperiments and showed that water must be regarded as a compound ofhydrogen and oxygen. ~General methods employed for the determination of the composition of acompound. ~ The composition of a compound may be determined by either oftwo general processes these are known as _analysis_ and _synthesis_. 1. _Analysis_ is the process of decomposing a compound into itsconstituents and determining what these constituents are. The analysisis _qualitative_ when it results in merely determining what elementscompose the compound; it is _quantitative_ when the exact percentage ofeach constituent is determined. Qualitative analysis must thereforeprecede quantitative analysis, for it must be known what elements, arein a compound before a method can be devised for determining exactly howmuch of each is present. 2. _Synthesis_ is the process of forming a compound from its constituentparts. It is therefore the reverse of analysis. Like analysis, it may beeither qualitative or quantitative. ~Application of these methods to the determination of the composition ofwater. ~ The determination of the composition of water is a matter ofgreat interest not only because of the importance of the compound butalso because the methods employed illustrate the general methods ofanalysis and synthesis. ~Methods based on analysis. ~ The methods based on analysis may be eitherqualitative or quantitative in character. [Illustration: Fig. 18] 1. _Qualitative analysis. _ As was stated in the study of oxygen, watermay be separated into its component parts by means of the electriccurrent. The form of apparatus ordinarily used for effecting thisanalysis is shown in Fig. 18. A platinum wire, to the end of which isattached a small piece of platinum foil (about 15 mm. By 25 mm. ), isfused through each of the tubes B and D, as shown in the figure. Thestopcocks at the ends of these tubes are opened and water, to which hasbeen added about one tenth of its volume of sulphuric acid, is pouredinto the tube A until the side tubes B and D are completelyfilled. The stopcocks are then closed. The platinum wires extending intothe tubes B and D are now connected with the wires leading from twoor three dichromate cells joined in series. The pieces of platinum foilwithin the tubes thus become the electrodes, and the current flows fromone to the other through the acidulated water. As soon as the currentpasses, bubbles of gas rise from each of the electrodes and collect inthe upper part of the tubes. The gas rising from the negative electrodeis found to be hydrogen, while that from the positive electrode isoxygen. It will be seen that the volume of the hydrogen is approximatelydouble that of the oxygen. Oxygen is more soluble in water thanhydrogen, and a very little of it is also lost by being converted intoozone and other substances. It has been found that when the necessarycorrections are made for the error due to these facts, the volume of thehydrogen is exactly double that of the oxygen. Fig. 19 illustrates a simpler form of apparatus, which may be used inplace of that shown in Fig. 18. A glass or porcelain dish is partiallyfilled with water to which has been added the proper amount of acid. Twotubes filled with the same liquid are inverted over the electrodes. Thegases resulting from the decomposition of the water collect in thetubes. [Illustration: Fig. 19] 2. _Quantitative analysis. _ The analysis just described is purelyqualitative and simply shows that water contains hydrogen and oxygen. Itdoes not prove the absence of other elements; indeed it does not provethat the hydrogen and oxygen are present in the proportion in which theyare liberated by the electric current. The method may be madequantitative, however, by weighing the water decomposed and also thehydrogen and oxygen obtained in its decomposition. If the combinedweights of the hydrogen and oxygen exactly equal the weight of the waterdecomposed, then it would be proved that the water consists of hydrogenand oxygen in the proportion in which they are liberated by the electriccurrent. This experiment is difficult to carry out, however, so that themore accurate methods based on synthesis are used. ~Methods based on synthesis. ~ Two steps are necessary to ascertain theexact composition of water by synthesis: (1) to show by qualitativesynthesis that water is formed by the union of oxygen with hydrogen; (2)to determine by quantitative synthesis in what proportion the twoelements unite to form water. The fact that water is formed by thecombination of oxygen with hydrogen was proved in the preceding chapter. The quantitative synthesis may be made as follows: [Illustration: Fig. 20] The combination of the two gases is brought about in a tube called aeudiometer. This is a graduated tube about 60 cm. Long and 2 cm. Wide, closed at one end (Fig. 20). Near the closed end two platinum wires arefused through the glass, the ends of the wires within the tube beingseparated by a space of 2 mm or 3 mm. The tube is entirely filled withmercury and inverted in a vessel of the same liquid. Pure hydrogen ispassed into the tube until it is about one fourth filled. The volume ofthe gas is then read off on the scale and reduced to standardconditions. Approximately an equal volume of pure oxygen is thenintroduced and the volume again read off and reduced to standardconditions. This gives the total volume of the two gases. From this thevolume of the oxygen introduced may be determined by subtracting fromit the volume of the hydrogen. The combination of the two gases is nowbrought about by connecting the two platinum wires with an inductioncoil and passing a spark from one wire to the other. Immediately aslight explosion occurs. The mercury in the tube is at first depressedbecause of the expansion of the gases due to the heat generated, but atonce rebounds, taking the place of the gases which have combined to formwater. The volume of the water in the liquid state is so small that itmay be disregarded in the calculations. In order that the temperature ofthe residual gas and the mercury may become uniform, the apparatus isallowed to stand for a few minutes. The volume of the gas is then readoff and reduced to standard conditions, so that it may be compared withthe volumes of the hydrogen and oxygen originally taken. The residualgas is then tested in order to ascertain whether it is hydrogen oroxygen, experiments having proved that it is never a mixture of the two. From the information thus obtained the composition of the water may becalculated. Thus, suppose the readings were as follows: Volume of hydrogen taken 20. 3 cc. Volume of hydrogen and oxygen 38. 7Volume of oxygen 18. 4Volume of gas left after combination has taken place (oxygen) 8. 3 The 20. 3 cc. Of hydrogen have combined with 18. 4 cc. Minus 8. 3 cc. (or10. 1 cc. ) of oxygen; or approximately 2 volumes of hydrogen havecombined with 1 of oxygen. Since oxygen is 15. 88 times as heavy ashydrogen, the proportion by weight in which the two gases combine is 1part of hydrogen to 7. 94 of oxygen. ~Precaution. ~ If the two gases are introduced into the eudiometer in theexact proportions in which they combine, after the combination has takenplace the liquid will rise and completely fill the tube. Under theseconditions, however, the tube is very likely to be broken by the suddenupward rush of the liquid. Hence in performing the experiment care istaken to introduce an excess of one of the gases. ~A more convenient form of eudiometer. ~ A form of eudiometer (Fig. 21)different from that shown on page 43 is sometimes used to avoid thecalculations necessary in reducing the volumes of the gases to the sameconditions of temperature and pressure in order to make comparisons. With this apparatus it is possible to take the readings of the volumesunder the same conditions of temperature and pressure, and thus comparethem directly. The apparatus (Fig. 21) is filled with mercury and thegases introduced into the tube A. The experiment is carried out as inthe preceding one, except that before taking the reading of the gasvolumes, mercury is either added to the tube B or withdrawn from it bymeans of the stopcock C, until it stands at exactly the same height inboth tubes. The gas inclosed in tube A is then under atmosphericpressure; and since but a few minutes are required for performing theexperiment, the conditions of temperature and pressure may be regardedas constant. Hence the volumes of the hydrogen and oxygen and of theresidual gas may be read off from the tube and directly compared. [Illustration: Fig. 21] ~Method used by Berzelius and Dumas. ~ The method used by theseinvestigators enables us to determine directly the proportion by weightin which the hydrogen and oxygen combine. Fig. 22 illustrates theapparatus used in making this determination. B is a glass tubecontaining copper oxide. C and D are glass tubes filled with calciumchloride, a substance which has great affinity for water. The tubes Band C, including their contents, are carefully weighed, and theapparatus connected as shown in the figure. A slow current of purehydrogen is then passed through A, and that part of the tube B whichcontains copper oxide is carefully heated. The hydrogen combines withthe oxygen present in the copper oxide to form water, which is absorbedby the calcium chloride in tube C. The calcium chloride in tube Dprevents any moisture entering tube C from the air. The operation iscontinued until an appreciable amount of water has been formed. Thetubes B and C are then weighed once more. The loss of weight in thetube B will exactly equal the weight of oxygen taken up from thecopper oxide in the formation of the water. The gain in weight in thetube C will exactly equal the weight of the water formed. Thedifference in these weights will of course equal the weight of thehydrogen present in the water formed. [Illustration: Fig. 22] ~Dumas' results. ~ The above method for the determination of thecomposition of water was first used by Berzelius in 1820. The work wasrepeated in 1843 by Dumas, the average of whose results is as follows: Weight of water formed 236. 36 g. Oxygen given up by the copper oxide 210. 04 ------Weight of hydrogen present in water 26. 32 According to this experiment the ratio of hydrogen to oxygen in water istherefore 26. 32 to 210. 04, or as l to 7. 98 ~Morley's results. ~ The American chemist Morley has recently determinedthe composition of water, extreme precautions being taken to use purematerials and to eliminate all sources of error. The hydrogen and oxygenwhich combined, as well as the water formed, were all accuratelyweighed. According to Morley's results, 1 part of hydrogen by weightcombines with 7. 94 parts of oxygen to form water. ~Comparison of results obtained. ~ From the above discussions it is easy tosee that it is by experiment alone that the composition of a compoundcan be determined. Different methods may lead to slightly differentresults. The more accurate the method chosen and the greater the skillwith which the experiment is carried out, the more accurate will be theresults. It is generally conceded by chemists that the results obtainedby Morley in reference to the composition of water are the most accurateones. In accordance with these results, then, _water must be regarded asa compound containing hydrogen and oxygen in the proportion of 1 part byweight of hydrogen to 7. 94 parts by weight of oxygen_. ~Relation between the volume of aqueous vapor and the volumes of thehydrogen and oxygen which combine to form it. ~ When the quantitativesynthesis of water is carried out in the eudiometer as described above, the water vapor formed by the union of the hydrogen and oxygen at oncecondenses. The volume of the resulting liquid is so small that it may bedisregarded in making the calculations. If, however, the experiment iscarried out at a temperature of 100° or above, the water-vapor formed isnot condensed and it thus becomes possible to compare the volume of thevapor with the volumes of hydrogen and oxygen which combined to form it. This can be accomplished by surrounding the arm A of the eudiometer(Fig. 23) with the tube B through which is passed the vapor obtainedby boiling some liquid which has a boiling point above 100°. In this wayit has been proved that 2 volumes of hydrogen and 1 volume of oxygencombine to form exactly 2 volumes of water vapor, the volumes all beingmeasured under the same conditions of temperature and pressure. It willbe noted that the relation between these volumes may be expressed bywhole numbers. The significance of this very important fact will bediscussed in a subsequent chapter. [Illustration: Fig. 23] ~Occurrence of water. ~ Water not only covers about three fourths of thesurface of the earth, and is present in the atmosphere in the form ofmoisture, but it is also a common constituent of the soil and rocks andof almost every form of animal and vegetable organism. The human body isnearly 70% water. This is derived not only from the water which we drinkbut also from the food which we eat, most of which contains a largepercentage of water. Thus potatoes contain about 78% of water, milk 85%, beef over 50%, apples 84%, tomatoes 94%. ~Impurities in water. ~ Chemically pure water contains only hydrogen andoxygen. Such a water never occurs in nature, however, for being a goodsolvent, it takes up certain substances from the rocks and soil withwhich it comes in contact. When such waters are evaporated thesesubstances are deposited in the form of a residue. Even rain water, which is the purest form occurring in nature, contains dust particlesand gases dissolved from the atmosphere. The foreign matter in water isof two kinds, namely, _mineral_, such as common salt and limestone, and_organic_, that is the products of animal and vegetable life. ~Mineral matter in water. ~ The amount and nature of the mineral matter present in different waters vary greatly, depending on the character of the rocks and soil with which the waters come in contact. The more common of the substances present are common salt and compounds of calcium, magnesium, and iron. One liter of the average river water contains about 175 mg. Of mineral matter. Water from deep wells naturally contains more mineral matter than river water, generally two or three times as much, while sea water contains as much as 35, 000 mg. To the liter. ~Effect of impurities on health. ~ The mineral matter in water does not, save in very exceptional cases, render the water injurious to the humansystem. In fact the presence of a certain amount of such matter isadvantageous, supplying the mineral constituents necessary for theformation of the solid tissues of the body. The presence of organicmatter, on the other hand, must always be regarded with suspicion. Thisorganic matter may consist not only of the products of animal andvegetable life but also of certain microscopic forms of living organismswhich are likely to accompany such products. Contagious diseases areknown to be due to the presence in the body of minute living organismsor germs. Each disease is caused by its own particular kind of germ. Through sewage these germs may find their way from persons afflictedwith disease into the water supply, and it is principally through thedrinking water that certain of these diseases, especially typhoid fever, are spread. It becomes of great importance, therefore, to be able todetect such matter when present in drinking water as well as to devisemethods whereby it can be removed or at least rendered harmless. ~Analysis of water. ~ The mineral analysis of a water is, as the name suggests, simply the determination of the mineral matter present. Sanitary analysis, on the other hand, is the determination of the organic matter present. The physical properties of a water give no conclusive evidence as to its purity, since a water may be unfit for drinking purposes and yet be perfectly clear and odorless. Neither can any reliance be placed on the simple methods often given for testing the purity of water. Only the trained chemist can carry out such methods of analysis as can be relied upon. [Illustration: Fig. 24] ~Purification of water. ~ Three general methods are used for thepurification of water, namely, _distillation_, _filtration_, and_boiling_. 1. _Distillation. _ The most effective way of purifying natural waters isby the process of distillation. This consists in boiling the water andcondensing the steam. Fig. 24 illustrates the process of distillation, as commonly conducted in the laboratory. Ordinary water is poured intothe flask A and boiled. The steam is conducted through the condenserB, which consists essentially of a narrow glass tube sealed within alarger one, the space between the two being filled with cold water, which is admitted at C and escapes at D. The inner tube is thus keptcool and the steam in passing through it is condensed. The water formedby the condensation of the steam collects in the receiver E and isknown as _distilled_ water. Such water is practically pure, since theimpurities are nonvolatile and remain in the flask A. ~Commercial distillation. ~ In preparing distilled water on a large scale, the steam is generated in a boiler or other metal container and condensed by passing it through a pipe made of metal, generally tin. This pipe is wound into a spiral and is surrounded by a current of cold water. Distilled water is used by the chemist in almost all of his work. It is also used in the manufacture of artificial ice and for drinking water. ~Fractional distillation. ~ In preparing distilled water, it is evident that if the natural water contains some substance which is volatile its vapor will pass over and be condensed with the steam, so that the distillate will not be pure water. Even such mixtures, however, may generally be separated by repeated distillation. Thus, if a mixture of water (boiling point 100°) and alcohol (boiling point 78°) is distilled, the alcohol, having the lower boiling point, tends to distill first, followed by the water. The separation of the two is not perfect, however, but may be made nearly so by repeated distillations. The process of separating a mixture of volatile substances by distillation is known as _fractional distillation_. 2. _Filtration. _ The process of distillation practically removes allnonvolatile foreign matter, mineral as well as organic. In purifyingwater for drinking purposes, however, it is only necessary to eliminatethe latter or to render it harmless. This is ordinarily done either byfiltration or boiling. In filtration the water is passed through somemedium which will retain the organic matter. Ordinary charcoal is aporous substance and will condense within its pores the organic matterin water if brought in contact with it. It is therefore well adapted tothe construction of filters. Such filters to be effective must be keptclean, since it is evident that the charcoal is useless after its poresare filled. A more effective type of filter is the Chamberlain-Pasteurfilter. In this the water is forced through a porous cylindrical cup, the pores being so minute as to strain out the organic matter. ~City filtration beds. ~ For purifying the water supply of cities, large filtration beds are prepared from sand and gravel, and the water is allowed to filter through these. Some of the impurities are strained out by the filter, while others are decomposed by the action of certain kinds of bacteria present in the sand. Fig. 25 shows a cross section of a portion of the filter used in purifying the water supply of Philadelphia. The water filters through the sand and gravel and passes into the porous pipe A, from which it is pumped into the city mains. The filters are covered to prevent the water from freezing in cold weather. [Illustration: Fig. 25] 3. _Boiling. _ A simpler and equally efficient method for purifying waterfor drinking purposes consists in boiling the water. It is the germs inwater that render it dangerous to health. These germs are living formsof matter. If the water is boiled, the germs are killed and the waterrendered safe. While these germs are destroyed by heat, cold has littleeffect upon them. Thus Dewar, in working with liquid hydrogen, exposedsome of these minute forms of life to the temperature of boilinghydrogen (-252°) without killing them. ~Self-purification of water. ~ It has long been known that watercontaminated with organic matter tends to purify itself when exposed tothe air. This is due to the fact that the water takes up a small amountof oxygen from the air, which gradually oxidizes the organic matterpresent in the water. While water is undoubtedly purified in this way, the method cannot be relied upon to purify a contaminated water so as torender it safe for drinking purposes. ~Physical properties. ~ Pure water is an odorless and tasteless liquid, colorless in thin layers, but having a bluish tinge when observedthrough a considerable thickness. It solidifies at 0° and boils at 100°under the normal pressure of one atmosphere. If the pressure isincreased, the boiling point is raised. When water is cooled it steadilycontracts until the temperature of 4° is reached: it then expands. Wateris remarkable for its ability to dissolve other substances, and is thebest solvent known. Solutions of solids in water are more frequentlyemployed in chemical work than are the solid substances, for chemicalaction between substances goes on more readily when they are in solutionthan it does when they are in the solid state. ~Chemical properties. ~ Water is a very stable substance, or, in otherwords, it does not undergo decomposition readily. To decompose it intoits elements by heat alone requires a very high temperature; at 2500°, for example, only about 5% of the entire amount is decomposed. Thoughvery stable towards heat, water can be decomposed in other ways, as bythe action of the electrical current or by certain metals. ~Heat of formation and heat of decomposition are equal. ~ The fact that a very high temperature is necessary to decompose water into hydrogen and oxygen is in accord with the fact that a great deal of heat is evolved by the union of hydrogen and oxygen; for it has been proved that the heat necessary to decompose a compound into its elements (heat of decomposition) is equal to the heat evolved in the formation of a compound from its elements (heat of formation). ~Water of crystallization. ~ When a solid is dissolved in water and theresulting solution is allowed to evaporate, the solid separates out, often in the form of crystals. It has been found that the crystals ofmany compounds, although perfectly dry, give up a definite amount ofwater when heated, the substance at the same time losing its crystallineform. Such water is called _water of crystallization_. This varies inamount with different compounds, but is perfectly definite for the samecompound. Thus, if a perfectly dry crystal of copper sulphate isstrongly heated in a tube, water is evolved and condenses on the sidesof the tube, the crystal crumbling to a light powder. The weight of thewater evolved is always equal to exactly 36. 07% of the weight of coppersulphate crystals heated. The water must therefore be in chemicalcombination with the substance composing the crystal; for if simplymixed with it or adhering to it, not only would the substance appearmoist but the amount present would undoubtedly vary. The combination, however, must be a very weak one, since the water is often expelled byeven a gentle heat. Indeed, in some cases the water is given up onsimple exposure to air. Such compounds are said to be _efflorescent_. Thus a crystal of sodium sulphate (Glauber's salt) on exposure to aircrumbles to a fine powder, owing to the escape of its water ofcrystallization. Other substances have just the opposite property: theyabsorb moisture when exposed to the air. For example, if a bit of drycalcium chloride is placed in moist air, in the course of a few hours itwill have absorbed sufficient moisture to dissolve it. Such substancesare said to be _deliquescent_. A deliquescent body serves as a gooddrying or _desiccating_ agent. We have already employed calcium chlorideas an agent for absorbing the moisture from hydrogen. Many substances, as for example quartz, form crystals which contain no water ofcrystallization. ~Mechanically inclosed water. ~ Water of crystallization must be carefully distinguished from water which is mechanically inclosed in a crystal and which can be removed by powdering the crystal and drying. Thus, when crystals of common salt are heated, the water inclosed in the crystal is changed into steam and bursts the crystal with a crackling sound. Such crystals are said to _decrepitate_. That this water is not combined is proved by the fact that the amount present varies and that it has all the properties of water. ~Uses of water. ~ The importance of water in its relation to life andcommerce is too well known to require comment. Its importance to thechemist has also been pointed out. It remains to call attention to thefact that it is used as a standard in many physical measurements. Thus0° and 100° on the centigrade scale are respectively the freezing andthe boiling points of water under normal pressure. The weight of 1 cc. Of water at its point of greatest density is the unit of weight in themetric system, namely, the gram. It is also taken as the unit for thedetermination of the density of liquids and solids as well as for themeasurement of amounts of heat. HYDROGEN DIOXIDE ~Composition. ~ As has been shown, 1 part by weight of hydrogen combineswith 7. 94 parts by weight of oxygen to form water. It is possible, however, to obtain a second compound of hydrogen and oxygen differingfrom water in composition in that 1 part by weight of hydrogen iscombined with 2 × 7. 94, or 15. 88 parts, of oxygen. This compound iscalled _hydrogen dioxide_ or _hydrogen peroxide_, the prefixes _di-_ and_per-_ signifying that it contains more oxygen than hydrogen oxide, which is the chemical name for water. ~Preparation. ~ Hydrogen dioxide cannot be prepared cheaply by the directunion of hydrogen and oxygen, and indirect methods must therefore beused. It is commonly prepared by the action of a solution of sulphuricacid on barium dioxide. The change which takes place may be indicated asfollows: sulphuric acid + barium dioxide = barium sulphate + hydrogen dioxide-------------- -------------- --------------- ---------------- hydrogen barium barium hydrogen sulphur oxygen sulphur oxygen oxygen oxygen In other words, the barium and hydrogen in the two compounds exchangeplaces. By this method a dilute solution of the dioxide in water isobtained. It is possible to separate the dioxide from the water byfractional distillation. This is attended with great difficulties, however, since the pure dioxide is explosive. The distillation iscarried on under diminished pressure so as to lower the boiling pointsas much as possible; otherwise the high temperature would decompose thedioxide. ~Properties. ~ Pure hydrogen dioxide is a colorless sirupy liquid having adensity of 1. 49. Its most characteristic property is the ease with whichit decomposes into water and oxygen. One part by weight of hydrogen iscapable of holding firmly only 7. 94 parts of oxygen. The additional 7. 94parts of oxygen present in hydrogen dioxide are therefore easilyevolved, the compound breaking down into water and oxygen. Thisdecomposition is attended by the generation of considerable heat. Indilute solution hydrogen dioxide is fairly stable, although such asolution should be kept in a dark, cool place, since both heat and lightaid in the decomposition of the dioxide. ~Uses. ~ Solutions of hydrogen dioxide are used largely as oxidizingagents. The solution sold by druggists contains 3% of the dioxide and isused in medicine as an antiseptic. Its use as an antiseptic depends uponits oxidizing properties. EXERCISES 1. Why does the chemist use distilled water in making solutions, ratherthan filtered water? 2. How could you determine the total amount of solid matter dissolved ina sample of water? 3. How could you determine whether a given sample of water is distilledwater? 4. How could the presence of air dissolved in water be detected? 5. How could the amount of water in a food such as bread or potato bedetermined? 6. Would ice frozen from impure water necessarily be free from diseasegerms? 7. Suppose that the maximum density of water were at 0° in place of 4°;what effect would this have on the formation of ice on bodies of water? 8. Is it possible for a substance to contain both mechanically inclosedwater and water of crystallization? 9. If steam is heated to 2000° and again cooled, has any chemical changetaken place in the steam? 10. Why is cold water passed into C instead of D (Fig. 24)? 11. Mention at least two advantages that a metal condenser has over aglass condenser. 12. Draw a diagram of the apparatus used in your laboratory forsupplying distilled water. 13. 20 cc. Of hydrogen and 7 cc. Of oxygen are placed in a eudiometerand the mixture exploded. (a) How many cubic centimeters of aqueousvapor are formed? (b) What gas and how much of it remains in excess? 14. (a) What weight of water can be formed by the combustion of 100 Lof hydrogen, measured under standard conditions? (b)What volume ofoxygen would be required in (a)? (c)What weight of potassiumchlorate is necessary to prepare this amount of oxygen? 15. What weight of oxygen is present in 1 kg. Of the ordinary hydrogendioxide solution? In the decomposition of this weight of the dioxideinto water and oxygen, what volume of oxygen (measured under standardconditions) is evolved? CHAPTER V THE ATOMIC THEORY ~Three fundamental laws of matter. ~ Before we can gain any very definiteidea in regard to the structure of matter, and the way in whichdifferent kinds of substances act chemically upon each other, it isnecessary to have clearly in view three fundamental laws of matter. These laws have been established by experiment, and any conception whichmay be formed concerning matter must therefore be in harmony with them. The laws are as follows: ~Law of conservation of matter. ~ This law has already been touched upon inthe introductory chapter, and needs no further discussion. It will berecalled that it may be stated thus: _Matter can neither be created nordestroyed, though it can be changed from one form into another. _ ~Law of definite composition. ~ In the earlier days of chemistry there wasmuch discussion as to whether the composition of a given compound isalways precisely the same or whether it is subject to some variation. Two Frenchmen, Berthollet and Proust, were the leaders in thisdiscussion, and a great deal of most useful experimenting was done todecide the question. Their experiments, as well as all succeeding ones, have shown that the composition of a pure chemical compound is alwaysexactly the same. Water obtained by melting pure ice, condensing steam, burning hydrogen in oxygen, has always 11. 18% hydrogen and 88. 82% oxygenin it. Red oxide of mercury, from whatever source it is obtained, contains 92. 6% mercury and 7. 4% oxygen. This truth is known as _the lawof definite composition_, and may be stated thus: _The composition of achemical compound never varies. _ ~Law of multiple proportion. ~ It has already been noted, however, thathydrogen and oxygen combine in two different ratios to form water andhydrogen dioxide respectively. It will be observed that this fact doesnot contradict the law of definite composition, for entirely differentsubstances are formed. These compounds differ from each other incomposition, but the composition of each one is always constant. Thisability of two elements to unite in more than one ratio is veryfrequently observed. Carbon and oxygen combine in two different ratios;nitrogen and oxygen combine to form as many as five distinct compounds, each with its own precise composition. In the first decade of the last century John Dalton, an Englishschool-teacher and philosopher, endeavored to find some rule which holdsbetween the ratios in which two given substances combine. His studiesbrought to light a very simple relation, which the following exampleswill make clear. In water the hydrogen and oxygen are combined in theratio of 1 part by weight of hydrogen to 7. 94 parts by weight of oxygen. In hydrogen dioxide the 1 part by weight of hydrogen is combined with15. 88 parts by weight of oxygen. The ratio between the amounts of oxygenwhich combine with the same amount of hydrogen to form water andhydrogen dioxide respectively is therefore 7. 94: 15. 88, or 1: 2. [Illustration: JOHN DALTON (English) (1766-1844) Developed the atomic theory; made many studies on the properties and thecomposition of gases. His book entitled "A New System of ChemicalPhilosophy" had a large influence on the development of chemistry] Similarly, the element iron combines with oxygen to form two oxides, oneof which is black and the other red. By analysis it has been shown thatthe former contains 1 part by weight of iron combined with 0. 286 partsby weight of oxygen, while the latter contains 1 part by weight of ironcombined with 0. 429 parts by weight of oxygen. Here again we find thatthe amounts of oxygen which combine with the same fixed amount of ironto form the two compounds are in the ratio of small whole numbers, viz. , 2:3. Many other examples of this simple relation might be given, since it hasbeen found to hold true in all cases where more than one compound is, formed from the same elements. Dalton's law of multiple proportionstates these facts as follows: _When any two elements, _ A _and_ B, _combine to form more than one compound, the amounts of_ B _which unitewith any fixed amount of_ A _bear the ratio of small whole numbers toeach other_. ~Hypothesis necessary to explain the laws of matter. ~ These threegeneralizations are called _laws_, because they express in conciselanguage truths which are found by careful experiment to hold good inall cases. They do not offer any explanation of the facts, but merelystate them. The human mind, however, does not rest content with the merebare facts, but seeks ever to learn the explanation of the facts. Asuggestion which is offered to explain such a set of facts is called an_hypothesis_. The suggestion which Dalton offered to explain the threelaws of matter, called the _atomic hypothesis_, was prompted by his viewof the constitution of matter, and it involves three distinctassumptions in regard to the nature of matter and chemical action. Dalton could not prove these assumptions to be true, but he saw that ifthey were true the laws of matter become very easy to understand. ~Dalton's atomic hypothesis. ~ The three assumptions which Dalton made inregard to the nature of matter, and which together constitute the atomichypothesis, are these: 1. All elements are made up of minute, independent particles whichDalton designated as _atoms_. 2. All atoms of the same element have equal masses; those of differentelements have different masses; in any change to which an atom issubjected its mass does not change. 3. When two or more elements unite to form a compound, the actionconsists in the union of a definite small number of atoms of eachelement to form a small particle of the compound. The smallest particlesof a given compound are therefore exactly alike in the number and kindsof atoms which they contain, and larger masses of the substances aresimply aggregations of these least particles. ~Molecules and atoms. ~ Dalton applied the name atom not only to the minuteparticles of the elements but also to the least particles of compounds. Later Avogadro, an Italian scientist, pointed out the fact that the twoare different, since the smallest particle of an element is a unit, while that of a compound must have at least two units in it. Hesuggested the name _molecule_ for the least particle of a compound whichcan exist, retaining the name _atom_ for the smallest particle of anelement. In accordance with this distinction, we may define the atom andthe molecule as follows: _An atom is the smallest particle of an elementwhich can exist. A molecule is the smallest particle of a compound whichcan exist. _ It will be shown in a subsequent chapter that sometimes twoor more atoms of the same element unite with each other to formmolecules of the element. While the term atom, therefore, is applicableonly to elements, the term molecule is applicable both to elements andcompounds. ~The atomic hypothesis and the laws of matter. ~ Supposing the atomichypothesis to be true, let us now see if it is in harmony with the lawsof matter. 1. _The atomic hypothesis and the law of conservation of matter. _ It isevident that if the atoms never change their masses in any change whichthey undergo, the total quantity of matter can never change and the lawof conservation of matter must follow. 2. _The atomic hypothesis and the law of definite composition. _According to the third supposition, when iron combines with sulphur theunion is between definite numbers of the two kinds of atoms. In thesimplest case one atom of the one element combines with one atom of theother. If the sulphur and the iron atoms never change their respectivemasses when they unite to form a molecule of iron sulphide, all ironsulphide molecules will have equal amounts of iron in them and also ofsulphur. Consequently any mass made up of iron sulphide molecules willhave the same fraction of iron by weight as do the individual ironsulphide molecules. Iron sulphide, from whatever source, will have thesame composition, which is in accordance with the law of definitecomposition. 3. _The atomic hypothesis and the law of multiple proportion. _ But thissimplest case may not always be the only one. Under other conditions oneatom of iron might combine with two of sulphur to form a molecule of asecond compound. In such a case the one atom of iron would be incombination with twice the mass of sulphur that is in the firstcompound, since the sulphur atoms all have equal masses. What is truefor one molecule will be true for any number of them; consequently whensuch quantities of these two compounds are selected as are found tocontain the same amount of iron, the one will contain twice as muchsulphur as the other. The combination between the atoms may of course take place in othersimple ratios. For example, two atoms of one element might combine withthree or with five of the other. In all such cases it is clear that thelaw of multiple proportion must hold true. For on selecting such numbersof the two kinds of molecules as have the same number of the one kind ofatoms, the numbers of the other kind of atoms will stand in some simpleratio to each other, and their weights will therefore stand in the samesimple ratio. ~Testing the hypothesis. ~ Efforts have been made to find compounds whichdo not conform to these laws, but all such attempts have resulted infailure. If such compounds should be found, the laws would be no longertrue, and the hypothesis of Dalton would cease to possess value. When anhypothesis has been tested in every way in which experiment can test it, and is still found to be in harmony with the facts in the case, it istermed a _theory_. We now speak of the atomic theory rather than of theatomic hypothesis. ~Value of a theory. ~ The value of a theory is twofold. It aids in theclear understanding of the laws of nature because it gives anintelligent idea as to why these laws should be in operation. A theory also leads to discoveries. It usually happens that in testing atheory much valuable work is done, and many new facts are discovered. Almost any theory in explaining given laws will involve a number ofconsequences apart from the laws it seeks to explain. Experiment willsoon show whether these facts are as the theory predicts they will be. Thus Dalton's atomic theory predicted many properties of gases whichexperiment has since verified. ~Atomic weights. ~ It would be of great advantage in the study of chemistryif we could determine the weights of the different kinds of atoms. It isevident that this cannot be done directly. They are so small that theycannot be seen even with a most powerful microscope. It is calculatedthat it would take 200, 000, 000 hydrogen atoms placed side by side tomake a row one centimeter long. No balance can weigh such minuteobjects. It is possible, however, to determine their relativeweights, --that is, how much heavier one is than another. _These relativeweights of the atoms are spoken of as the atomic weights of theelements. _ If elements were able to combine in only one way, --one atom of one withone atom of another, --the problem of determining the atomic weightswould be very simple. We should merely have to take some one convenientelement as a standard, and find by experiment how much of each otherelement would combine with a fixed weight of it. The ratios thus foundwould be the same ratios as those between the atoms of the elements, andthus we should have their relative atomic weights. The law of multipleproportion calls attention to the fact that the atoms combine in otherratios than 1: 1, and there is no direct way of telling which one, ifany, of the several compounds in a given case is the one consisting of asingle atom of each element. If some way were to be found of telling how much heavier the entiremolecule of a compound is than the atom chosen as a standard, --that is, of determining the molecular weights of compounds, --the problem could besolved, though its solution would not be an entirely simple matter. There are ways of determining the molecular weights of compounds, andthere are other experiments which throw light directly upon the relativeweights of the atoms. These methods cannot be described until the factsupon which they rest have been studied. It will be sufficient for thepresent to assume that these methods are trustworthy. ~Standard for atomic weights. ~ Since the atomic weights are merelyrelative to some one element chosen as a standard, it is evident thatany one of the elements may serve as this standard and that anyconvenient value may be assigned to its atom. At one time oxygen wastaken as this standard, with the value 100, and the atomic weights ofthe other elements were expressed in terms of this standard. It wouldseem more rational to take the element of smallest atomic weight as thestandard and give it unit value; accordingly hydrogen was taken as thestandard with an atomic weight of 1. Very recently, however, this unithas been replaced by oxygen, with an atomic weight of 16. ~Why oxygen is chosen as the standard for atomic weights. ~ In thedetermination of the atomic weight of an element it is necessary to findthe weight of the element which combines with a definite weight ofanother element, preferably the element chosen as the standard. Sinceoxygen combines with the elements far more readily than does hydrogen toform definite compounds, it is far better adapted for the standardelement, and has accordingly replaced hydrogen as the standard. Anydefinite value might be given to the weight of the oxygen atom. Inassigning a value to it, however, it is convenient to choose a wholenumber, and as small a number as possible without making the atomicweight of any other element less than unity. For these reasons thenumber 16 has been chosen as the atomic weight of oxygen. This makesthe atomic weight of hydrogen equal to 1. 008, so that there is butlittle difference between taking oxygen as 16 and hydrogen as 1 for theunit. The atomic weights of the elements are given in the Appendix. EXERCISES 1. Two compounds were found to have the following compositions: (a)oxygen = 69. 53%, nitrogen = 30. 47%; (b) oxygen = 53. 27%, nitrogen =46. 73%. Show that the law of multiple proportion holds in this case. 2. Two compounds were found to have the following compositions: (a)oxygen = 43. 64%, phosphorus = 56. 36%; (b) oxygen = 56. 35%, phosphorus= 43. 65%. Show that the law of multiple proportion holds in this case. 3. Why did Dalton assume that all the atoms of a given element have thesame weight? CHAPTER VI CHEMICAL EQUATIONS AND CALCULATIONS ~Formulas. ~ Since the molecule of any chemical compound consists of adefinite number of atoms, and this number never changes withoutdestroying the identity of the compound, it is very convenient torepresent the composition of a compound by indicating the composition ofits molecules. This can be done very easily by using the symbols of theatoms to indicate the number and the kind of the atoms which constitutethe molecule. HgO will in this way represent mercuric oxide, a moleculeof which has been found to contain 1 atom each of mercury and oxygen. H_{2}O will represent water, the molecules of which consist of 1 atom ofoxygen and 2 of hydrogen, the subscript figure indicating the number ofthe atoms of the element whose symbol precedes it. H_{2}SO_{4} willstand for sulphuric acid, the molecules of which contain 2 atoms ofhydrogen, 1 of sulphur, and 4 of oxygen. The combination of symbolswhich represents the molecule of a substance is called its _formula_. ~Equations. ~ When a given substance undergoes a chemical change it ispossible to represent this change by the use of such symbols andformulas. In a former chapter it was shown that mercuric oxidedecomposes when heated to form mercury and oxygen. This may be expressedvery briefly in the form of the equation (1) HgO = Hg + O. When water is electrolyzed two new substances, hydrogen and oxygen, areformed from it. This statement in the form of an equation is (2) H_{2}O = 2H + O. The coefficient before the symbol for hydrogen indicates that a singlemolecule of water yields two atoms of hydrogen on decomposition. In like manner the combination of sulphur with iron is expressed by theequation (3) Fe + S = FeS. The decomposition of potassium chlorate by heat takes place asrepresented by the equation (4) KClO_{3} = KCl + 3O. ~Reading of equations. ~ Since equations are simply a kind of shorthand wayof indicating chemical changes which occur under certain conditions, inreading an equation the full statement for which it stands should begiven. Equation (1) should be read, "Mercuric oxide when heated givesmercury and oxygen"; equation (2) is equivalent to the statement, "Whenelectrolyzed, water produces hydrogen and oxygen"; equation (3), "Whenheated together iron and sulphur unite to form iron sulphide"; equation(4), "Potassium chlorate when heated yields potassium chloride andoxygen. " ~Knowledge required for writing equations. ~ In order to write suchequations correctly, a considerable amount of exact knowledge isrequired. Thus, in equation (1) the fact that red oxide of mercury hasthe composition represented by the formula HgO, that it is decomposed byheat, that in this decomposition mercury and oxygen are formed and noother products, --all these facts must be ascertained by exact experimentbefore the equation can be written. An equation expressing these factswill then have much value. Having obtained an equation describing the conduct of mercuric oxide onbeing heated, it will not do to assume that other oxides will behave inlike manner. Iron oxide (FeO) resembles mercuric oxide in many respects, but it undergoes no change at all when heated. Manganese dioxide, theblack substance used in the preparation of oxygen, has the formulaMnO_{2}. When this substance is heated oxygen is set free, but the metalmanganese is not liberated; instead, a different oxide of manganesecontaining less oxygen is produced. The equation representing thereaction is 3MnO_{2} = Mn_{3}O_{4} + 2O. ~Classes of reactions. ~ When a chemical change takes place in a substancethe substance is said to undergo a reaction. Although a great manydifferent reactions will be met in the study of chemistry, they may allbe grouped under the following heads. 1. _Addition. _ This is the simplest kind of chemical action. It consistsin the union of two or more substances to produce a new substance. Thecombination of iron with sulphur is an example: Fe + S = FeS. 2. _Decomposition. _ This is the reverse of addition, the substanceundergoing reaction being parted into its constituents. Thedecomposition of mercuric oxide is an example: HgO = Hg + O. 3. _Substitution. _ It is sometimes possible for an element in the freestate to act upon a compound in such a way that it takes the place ofone of the elements of the compound, liberating it in turn. In the studyof the element hydrogen it was pointed out that hydrogen is mostconveniently prepared by the action of sulphuric or hydrochloric acidupon zinc. When sulphuric acid is used a substance called zinc sulphate, having the composition represented by the formula ZnSO_{4}, is formedtogether with hydrogen. The equation is Zn + H_{2}SO_{4} = ZnSO_{4} + 2H. When hydrochloric acid is used zinc chloride and hydrogen are theproducts of reaction: Zn + 2HCl = ZnCl_{2} + 2H. When iron is used in place of zinc the equation is Fe + H_{2}SO_{4} = FeSO_{4} + 2H. These reactions are quite similar, as is apparent from an examination ofthe equations. In each case 1 atom of the metal replaces 2 atoms ofhydrogen in the acid, and the hydrogen escapes as a gas. When an elementin the free state, such as the zinc in the equations just given, takesthe place of some one element in a compound, setting it free fromchemical combination, the act is called _substitution_. Other reactions illustrating substitution are the action of sodium onwater, Na + H_{2}O = NaOH + H; and the action of heated iron upon water, 3Fe + 4H_{2}O = Fe_{3}O_{4} + 8H. 4. _Double decomposition. _ When barium dioxide (BaO_{2}) is treated withsulphuric acid two compounds are formed, namely, hydrogen dioxide(H_{2}O_{2}) and barium sulphate (BaSO_{4}). The equation is BaO_{2} + H_{2}SO_{4} = BaSO_{4} + H_{2}O_{2}. In this reaction it will be seen that the two elements barium andhydrogen simply exchange places. Such a reaction is called a _doubledecomposition_. We shall meet with many examples of this kind ofchemical reactions. ~Chemical equations are quantitative. ~ The use of symbols and formulas inexpressing chemical changes has another great advantage. Thus, accordingto the equation H_{2}O = 2H + O, 1 molecule of water is decomposed into 2 atoms of hydrogen and 1 atom ofoxygen. But, as we have seen, the relative weights of the atoms areknown, that of hydrogen being 1. 008, while that of oxygen is 16. Themolecule of water, being composed of 2 atoms of hydrogen and 1 atom ofoxygen, must therefore weigh relatively 2. 016 + 16, or 18. 016. Theamount of hydrogen in this molecule must be 2. 016/18. 016, or 11. 18% ofthe whole, while the amount of oxygen must be 16/18. 018, or 88. 82% ofthe whole. Now, since any definite quantity of water is simply the sumof a great many molecules of water, it is plain that the fractionsrepresenting the relative amounts of hydrogen and oxygen present in amolecule must likewise express the relative amounts of hydrogen andoxygen present in any quantity of water. Thus, for example, in 20 g. Ofwater there are 2. 016/18. 016 × 20, or 2. 238 g. Of hydrogen, and16/18. 016 × 20, or 17. 762 g. Of oxygen. These results in reference tothe composition of water of course agree exactly with the facts obtainedby the experiments described in the chapter on water, for it is becauseof those experiments that the values 1. 008 and 16 are given to hydrogenand oxygen respectively. It is often easier to make calculations of this kind in the form of aproportion rather than by fractions. Since the molecule of water andthe two atoms of hydrogen which it contains have the ratio by weight of18. 016: 2. 016, any mass of water has the same ratio between its totalweight and the weight of the hydrogen in it. Hence, to find the numberof grams (x) of hydrogen in 20 g. Of water, we have the proportion 18. 016 : 2. 016 :: 20 g. : x (grams of hydrogen). Solving for x, we get 2. 238 for the number of grams of hydrogen. Similarly, to find the amount (x) of oxygen present in the 20 g. Ofwater, we have the proportion 18. 016 : 16 :: 20 : x from which we find that x = 17. 762 g. Again, suppose we wish to find what weight of oxygen can be obtainedfrom 15 g. Of mercuric oxide. The equation representing thedecomposition of mercuric oxide is HgO = Hg + O. The relative weights of the mercury and oxygen atoms are respectively200 and 16. The relative weight of the mercuric oxide molecule musttherefore be the sum of these, or 216. The molecule of mercuric oxideand the atom of oxygen which it contains have the ratio 216: 16. Thissame ratio must therefore hold between the weight of any given quantityof mercuric oxide and that of the oxygen which it contains. Hence, tofind the weight of oxygen in 15 g. Of mercuric oxide, we have theproportion 216 : 16 :: 15 : x (grams of oxygen). On the other hand, suppose we wish to prepare, say, 20 g. Of oxygen. Theproblem is to find out what weight of mercuric oxide will yield 20 g. Ofoxygen. The following proportion evidently holds 216 : 16 :: x (grams of mercuric oxide) : 20; from which we get x = 270. In the preparation of hydrogen by the action of sulphuric acid uponzinc, according to the equation, Zn + H_{2}SO_{4} = ZnSO_{4} + 2 H, suppose that 50 g. Of zinc are available; let it be required tocalculate the weight of hydrogen which can be obtained. It will be seenthat 1 atom of zinc will liberate 2 atoms of hydrogen. The ratio byweight of a zinc to an hydrogen atom is 65. 4: 1. 008; of 1 zinc atom to 2hydrogen atoms, 65. 4: 2. 016. Zinc and hydrogen will be related in thisreaction in this same ratio, however many atoms of zinc are concerned. Consequently in the proportion 65. 4 : 2. 016 :: 50 : x, x will be the weight of hydrogen set free by 50 g. Of zinc. The weightof zinc sulphate produced at the same time can be found from theproportion 65. 4 : 161. 46 :: 50 : x; where 161. 46 is the molecular weight of the zinc sulphate, and x theweight of zinc sulphate formed. In like manner, the weight of sulphuricacid used up can be calculated from the proportion 65. 4 : 98. 076 :: 50 : x. These simple calculations are possible because the symbols and formulasin the equations represent the relative weights of the substancesconcerned in a chemical reaction. When once the relative weights of theatoms have been determined, and it has been agreed to allow the symbolsto stand for these relative weights, an equation or formula making useof the symbols becomes a statement of a definite numerical fact, andcalculations can be based on it. ~Chemical equations not algebraic. ~ Although chemical equations arequantitative, it must be clearly understood that they are not algebraic. A glance at the equations 7 + 4 = 11, 8 + 5 = 9 + 4 will show at once that they are true. The equations HgO = Hg + O, FeO = Fe + O are equally true in an algebraic sense, but experiment shows that onlythe first is true chemically, for iron oxide (FeO) cannot be directlydecomposed into iron and oxygen. Only such equations as have been foundby careful experiment to express a real chemical transformation, trueboth for the kinds of substances as well as for the weights, have anyvalue. _Chemical formulas and equations, therefore, are a concise way ofrepresenting qualitatively and quantitatively facts which have beenfound by experiment to be true in reference to the composition ofsubstances and the changes which they undergo. _ ~Formulas representing water of crystallization. ~ An examination ofsubstances containing water of crystallization has shown that in everycase the water is present in such proportion by weight as can readily berepresented by a formula. For example, copper sulphate (CuSO_{4}) andwater combine in the ratio of 1 molecule of the sulphate to 5 of water;calcium sulphate (CaSO_{4}) and water combine in the ratio 1: 2 to formgypsum. These facts are expressed by writing the formulas for the twosubstances with a period between them. Thus the formula for crystallizedcopper sulphate is CuSO_{4}·5H_{2}O; that of gypsum is CaSO_{4}·2H_{2}O. ~Heat of reaction. ~ Attention has frequently been directed to the factthat chemical changes are usually accompanied by heat changes. Ingeneral it has been found that in every chemical action heat is eitherabsorbed or given off. By adopting a suitable unit for the measurementof heat, the heat change during a chemical reaction can be expressed inthe equation for the reaction. Heat cannot be measured by the use of a thermometer alone, since thethermometer measures the intensity of heat, not its quantity. Theeasiest way to measure a quantity of heat is to note how warm it willmake a definite amount of a given substance chosen as a standard. Waterhas been chosen as the standard, and the unit of heat is called a_calorie. A calorie is defined as the amount of heat required to raisethe temperature of one gram of water one degree. _ By means of this unit it is easy to indicate the heat changes in a givenchemical reaction. The equation 2H + O = H_{2}O + 68, 300 cal. means that when 2. 016 g. Of hydrogen combine with 16 g. Of oxygen, 18. 016 g. Of water are formed and 68, 300 cal. Are set free. C + 2S = CS_{2} - 19, 000 cal. means that an expenditure of 19, 000 cal. Is required to cause 12 g. Ofcarbon to unite with 64. 12 g. Of sulphur to form 76. 12 g. Of carbondisulphide. In these equations it will be noted that the symbols standfor as many grams of the substance as there are units in the weights ofthe atoms represented by the symbols. This is always understood to bethe case in equations where the heat of reaction is given. ~Conditions of a chemical action are not indicated by equations. ~Equations do not tell the conditions under which a reaction will takeplace. The equation HgO = Hg + O does not tell us that it is necessary to keep the mercuric oxide at ahigh temperature in order that the decomposition may go on. The equation Zn + 2HCl = ZnCl_{2} + 2H in no way indicates the fact that the hydrochloric acid must bedissolved in water before it will act upon the zinc. From the equation H + Cl = HCl it would not be suspected that the two gases hydrogen and chlorine willunite instantly in the sunlight, but will stand mixed in the dark a longtime without change. It will therefore be necessary to pay muchattention to the details of the conditions under which a given reactionoccurs, as well as to the expression of the reaction in the form of anequation. EXERCISES 1. Calculate the percentage composition of the following substances:(a) mercuric oxide; (b) potassium chlorate; (c) hydrochloric acid;(d) sulphuric acid. Compare the results obtained with the compositionsas given in Chapters II and III. 2. Determine the percentage of copper, sulphur, oxygen, and water incopper sulphate crystals. What weight of water can be obtained from 150g. Of this substance? 3. What weight of zinc can be dissolved in 10 g. Of sulphuric acid? Howmuch zinc sulphate will be formed? 4. How many liters of hydrogen measured under standard conditions can beobtained from the action of 8 g. Of iron on 10 g. Of sulphuric acid? Howmuch iron sulphate (FeSO_{4}) will be formed? 5. 10 g. Of zinc were used in the preparation of hydrogen; what weightof iron will be required to prepare an equal volume? 6. How many grams of barium dioxide will be required to prepare 1 kg. Ofcommon hydrogen dioxide solution? What weight of barium sulphate will beformed at the same time? 7. What weight of the compound Mn_{3}O_{4} will be formed by stronglyheating 25 g. Of manganese dioxide? What volume of oxygen will be givenoff at the same time, measured under standard conditions? 8. (a) What is the weight of 100 l. Of hydrogen measured in alaboratory in which the temperature is 20° and pressure 750 mm. ? (b)What weight of sulphuric acid is necessary to prepare this amount ofhydrogen? (c) The density of sulphuric acid is 1. 84. Express the acidrequired in (b) in cubic centimeters. 9. What weight of potassium chlorate is necessary to furnish sufficientoxygen to fill four 200 cc. Bottles in your laboratory (the gas to becollected over water)? CHAPTER VII NITROGEN AND THE RARE ELEMENTS: ARGON, HELIUM, NEON, KRYPTON, XENON ~Historical. ~ Nitrogen was discovered by the English chemist Rutherford in1772. A little later Scheele showed it to be a constituent of air, andLavoisier gave it the name _azote_, signifying that it would not supportlife. The name _nitrogen_ was afterwards given it because of itspresence in saltpeter or niter. The term azote and symbol Az are stillretained by the French chemists. ~Occurrence. ~ Air is composed principally of oxygen and nitrogen in thefree state, about 78 parts by volume out of every 100 parts beingnitrogen. Nitrogen also occurs in nature in the form of potassiumnitrate (KNO_{3})--commonly called saltpeter or niter--as well as insodium nitrate (NaNO_{3}). Nitrogen is also an essential constituent ofall living organisms; for example, the human body contains about 2. 4% ofnitrogen. ~Preparation from air. ~ Nitrogen can be prepared from air by the action ofsome substance which will combine with the oxygen, leaving the nitrogenfree. Such a substance must be chosen, however, as will combine with theoxygen to form a product which is not a gas, and which can be readilyseparated from the nitrogen. The substances most commonly used for thispurpose are phosphorus and copper. 1. _By the action of phosphorus. _ The method used for the preparation ofnitrogen by the action of phosphorus is as follows: The phosphorus is placed in a little porcelain dish, supported on a corkand floated on water (Fig. 26). It is then ignited by contact with a hotwire, and immediately a bell jar or bottle is brought over it so as toconfine a portion of the air. The phosphorus combines with the oxygen toform an oxide of phosphorus, known as phosphorus pentoxide. This is awhite solid which floats about in the bell jar, but in a short time itis all absorbed by the water, leaving the nitrogen. The withdrawal ofthe oxygen is indicated by the rising of the water in the bell jar. [Illustration: Fig. 26] 2. _By the action of copper. _ The oxygen present in the air may also beremoved by passing air slowly through a heated tube containing copper. The copper combines with the oxygen to form copper oxide, which is asolid. The nitrogen passes on and may be collected over water. ~Nitrogen obtained from air is not pure. ~ Inasmuch as air, in addition to oxygen and nitrogen, contains small amounts of other gases, and since the phosphorus as well as the copper removes only the oxygen, it is evident that the nitrogen obtained by these methods is never quite pure. About 1% of the product is composed of other gases, from which it is very difficult to separate the nitrogen. The impure nitrogen so obtained may, however, be used for a study of most of the properties of nitrogen, since these are not materially affected by the presence of the other gases. ~Preparation from compounds of nitrogen. ~ Pure nitrogen may be obtainedfrom certain compounds of the element. Thus, if heat is applied to thecompound ammonium nitrite (NH_{4}NO_{2}), the change represented in thefollowing equation takes place: NH_{4}NO_{2} = 2H_{2}O + 2N. ~Physical properties. ~ Nitrogen is similar to oxygen and hydrogen in thatit is a colorless, odorless, and tasteless gas. One liter of nitrogenweighs 1. 2501 g. It is almost insoluble in water. It can be obtained inthe form of a colorless liquid having a boiling point of -195° atordinary pressure. At -214° it solidifies. ~Chemical properties. ~ Nitrogen is characterized by its inertness. It isneither combustible nor a supporter of combustion. At ordinarytemperatures it will not combine directly with any of the elementsexcept under rare conditions. At higher temperatures it combines withmagnesium, lithium, titanium, and a number of other elements. Thecompounds formed are called _nitrides_, just as compounds of an elementwith oxygen are called _oxides_. When it is mixed with oxygen andsubjected to the action of electric sparks, the two gases slowly combineforming oxides of nitrogen. A mixture of nitrogen and hydrogen whentreated similarly forms ammonia, a gaseous compound of nitrogen andhydrogen. Since we are constantly inhaling nitrogen, it is evident thatit is not poisonous. Nevertheless life would be impossible in anatmosphere of pure nitrogen on account of the exclusion of the necessaryoxygen. ~Argon, helium, neon, krypton, xenon. ~ These are all rare elements occurring in the air in very small quantities. Argon, discovered in 1894, was the first one obtained. Lord Rayleigh, an English scientist, while engaged in determining the exact weights of various gases, observed that the nitrogen obtained from the air is slightly heavier than pure nitrogen obtained from its compounds. After repeating his experiments many times, always with the same results, Rayleigh finally concluded that the nitrogen which he had obtained from the air was not pure, but was mixed with a small amount of some unknown gas, the density of which is greater than that of nitrogen. Acting on this assumption, Rayleigh, together with the English chemist Ramsay, attempted to separate the nitrogen from the unknown gas. Knowing that nitrogen would combine with magnesium, they passed the nitrogen obtained from the air and freed from all known substances through tubes containing magnesium heated to the necessary temperature. After repeating this operation, they finally succeeded in obtaining from the atmospheric nitrogen a small volume of gas which would not combine with magnesium and hence could not be nitrogen. This proved to be a new element, to which they gave the name _argon_. As predicted, this new element was found to be heavier than nitrogen, its density as compared with hydrogen as a standard being approximately 20, that of nitrogen being only 14. About 1% of the atmospheric nitrogen proved to be argon. The new element is characterized by having no affinity for other elements. Even under the most favorable conditions it has not been made to combine with any other element. On this account it was given the name argon, signifying lazy or idle. Like nitrogen, it is colorless, odorless, and tasteless. It has been liquefied and solidified. Its boiling point is -187°. Helium was first found in the gases expelled from certain minerals by heating. Through the agency of the spectroscope it had been known to exist in the sun long before its presence on the earth had been demonstrated, --a fact suggested by the name helium, signifying the sun. Its existence in traces in the atmosphere has also been proven. It was first liquefied by Onnes in July, 1908. Its boiling point, namely -269°, is the lowest temperature yet reached. The remaining elements of this group--neon, krypton, and xenon--have been obtained from liquid air. When liquid air is allowed to boil, the constituents which are the most difficult to liquefy, and which therefore have the lowest boiling points, vaporize first, followed by the others in the order of their boiling points. It is possible in this way to make at least a partial separation of the air into its constituents, and Ramsay thus succeeded in obtaining from liquid air not only the known constituents, including argon and helium, but also the new elements, neon, krypton, and xenon. These elements, as well as helium, all proved to be similar to argon in that they are without chemical activity, apparently forming no compounds whatever. The percentages present in the air are very small. The names, neon, krypton, xenon, signify respectively, new, hidden, stranger. EXERCISES 1. How could you distinguish between oxygen, hydrogen, and nitrogen? 2. Calculate the relative weights of nitrogen and oxygen; of nitrogenand hydrogen. 3. In the preparation of nitrogen from the air, how would hydrogen do asa substance for the removal of the oxygen? 4. What weight of nitrogen can be obtained from 10 l. Of air measuredunder the conditions of temperature and pressure which prevail in yourlaboratory? 5. How many grams of ammonium nitrite are necessary in the preparationof 20 l. Of nitrogen measured over water under the conditions oftemperature and pressure which prevail in your laboratory? 6. If 10 l. Of air, measured under standard conditions, is passed over100 g. Of hot copper, how much will the copper gain in weight? [Illustration: WILLIAM RAMSAY (Scotch) (1855-) Has made many studies in the physical properties of substances;discovered helium; together with Lord Rayleigh and others he discoveredargon, krypton, xenon, and neon; has contributed largely to theknowledge of radio-active substances, showing that radium graduallygives rise to helium; professor at University College, London] CHAPTER VIII THE ATMOSPHERE ~Atmosphere and air. ~ The term _atmosphere_ is applied to the gaseousenvelope surrounding the earth. The term _air_ is generally applied to alimited portion of this envelope, although the two words are often usedinterchangeably. Many references have already been made to thecomposition and properties of the atmosphere. These statements must nowbe collected and discussed somewhat more in detail. ~Air formerly regarded as an element. ~ Like water, air was at firstregarded as elementary in character. Near the close of the eighteenthcentury Scheele, Priestley, and Lavoisier showed by their experimentsthat it is a mixture of at least two gases, --those which we now calloxygen and nitrogen. By burning substances in an inclosed volume of airand noting the contraction in volume due to the removal of the oxygen, they were able to determine with some accuracy the relative volumes ofoxygen and nitrogen present in the air. ~The constituents of the atmosphere. ~ The constituents of the atmospheremay be divided into two general groups: those which are essential tolife and those which are not essential. 1. _Constituents essential to life. _ In addition to oxygen and nitrogenat least two other substances, namely, carbon dioxide and water vapor, must be present in the atmosphere in order that life may exist. Theformer of these is a gaseous compound of carbon and oxygen having theformula CO_{2}. Its properties will be discussed in detail in thechapter on the compounds of carbon. Its presence in the air may be shownby causing the air to bubble through a solution of calcium hydroxide(Ca(OH)_{2}), commonly called lime water. The carbon dioxide combineswith the calcium hydroxide in accordance with the following equation: Ca(OH)_{2} + CO_{2} = CaCO_{3} + H_{2}O. The resulting calcium carbonate (CaCO_{3}) is insoluble in water andseparates in the form of a white powder, which causes the solution toappear milky. The presence of water vapor is readily shown by its condensation on coldobjects as well as by the fact that a bit of calcium chloride whenexposed to the air becomes moist, and may even dissolve in the waterabsorbed from the air. 2. _Constituents not essential to life. _ In addition to the essentialconstituents, the air contains small percentages of various other gases, the presence of which so far as is known is not essential to life. Thislist includes the rare elements, argon, helium, neon, krypton, andxenon; also hydrogen, ammonia, hydrogen dioxide, and probably ozone. Certain minute forms of life (germs) are also present, the decay oforganic matter being due to their presence. ~Function of each of the essential constituents. ~ (1) The oxygen directly supports life through respiration. (2) The nitrogen, on account of its inactivity, serves to dilute the oxygen, and while contrary to the older views, it is possible that life might continue to exist in the absence of the atmospheric nitrogen, yet the conditions of life would be entirely changed. Moreover, nitrogen is an essential constituent of all animal and plant life. It was formerly supposed that neither animals nor plants could assimilate the free nitrogen, but it has been shown recently that the plants of at least one natural order, the Leguminosę, to which belong the beans, peas, and clover, have the power of directly assimilating the free nitrogen from the atmosphere. This is accomplished through the agency of groups of bacteria, which form colonies in little tubercles on the roots of the plants. These bacteria probably assist in the absorption of nitrogen by changing the free nitrogen into compounds which can be assimilated by the plant. Fig. 27 shows the tubercles on the roots of a variety of bean. (3) The presence of water vapor in the air is necessary to prevent excessive evaporation from both plants and animals. (4) Carbon dioxide is an essential plant food. [Illustration: Fig. 27] ~The quantitative analysis of air. ~ A number of different methods havebeen devised for the determination of the percentages of theconstituents present in the atmosphere. Among these are the following. 1. _Determination of oxygen. _ (1) The oxygen is withdrawn from ameasured volume of air inclosed in a tube, by means of phosphorus. To make the determination, a graduated tube is filled with water and inverted in a vessel of water. Air is introduced into the tube until it is partially filled with the gas. The volume of the inclosed air is carefully noted and reduced to standard conditions. A small piece of phosphorus is attached to a wire and brought within the tube as shown in Fig. 28. After a few hours the oxygen in the inclosed air will have combined with the phosphorus, the water rising to take its place. The phosphorus is removed and the volume is again noted and reduced to standard conditions. The contraction in the volume of the air is equal to the volume of oxygen absorbed. [Illustration: Fig. 28] (2) The oxygen may also be estimated by passing a measured volume of airthrough a tube containing copper heated to a high temperature. Theoxygen in the air combines with the copper to form copper oxide (CuO). Hence the increase in the weight of the copper equals the weight of theoxygen in the volume of air taken. (3) A more accurate method is the following. A eudiometer tube is filledwith mercury and inverted in a vessel of the same liquid. A convenientamount of air is then introduced into the tube and its volume accuratelynoted. There is then introduced more than sufficient hydrogen to combinewith the oxygen present in the inclosed air, and the volume is againaccurately noted. The mixture is then exploded by an electric spark, andthe volume is once more taken. By subtracting this volume from the totalvolume of the air and hydrogen there is obtained the contraction involume due to the union of the oxygen and hydrogen. The volume occupiedby the water formed by the union of the two gases is so small that itmay be disregarded in the calculation. Since oxygen and hydrogen combinein the ratio 1: 2 by volume, it is evident that the contraction involume due to the combination is equal to the volume occupied by theoxygen in the air contained in the tube, plus twice this volume ofhydrogen. In other words, one third of the total contraction is equal tothe volume occupied by the oxygen in the inclosed air. The followingexample will make this clear: Volume of air in tube 50. 0 cc. Volume after introducing hydrogen 80. 0Volume after combination of oxygen and hydrogen 48. 5Contraction in volume due to combination (80 cc. -48. 5 cc. ) 31. 5Volume of oxygen in 50 cc. Of air (1/3 of 31. 5) 10. 5 All these methods agree in showing that 100 volumes of dry air containapproximately 21 volumes of oxygen. 2. _Determination of nitrogen. _ If the gas left after the removal ofoxygen from a portion of air is passed over heated magnesium, thenitrogen is withdrawn, argon and the other rare elements being left. Itmay thus be shown that of the 79 volumes of gas left after the removalof the oxygen from 100 volumes of air, approximately 78 are nitrogen and0. 93 argon. The other elements are present in such small quantities thatthey may be neglected. 3. _Determination of carbon dioxide. _ The percentage of carbon dioxidein any given volume of air may be determined by passing the air overcalcium hydroxide or some other compound which will combine with thecarbon dioxide. The increase in the weight of the hydroxide equals theweight of the carbon dioxide absorbed. The amount present in the opennormal air is from 3 to 4 parts by volume in 10, 000 volumes of air, orabout 0. 04%. 4. _Determination of water vapor. _ The water vapor present in a givenvolume of air may be determined by passing the air over calcium chloride(or some other compound which has a strong affinity for water), andnoting the increase in the weight of the chloride. The amount presentvaries not only with the locality, but there is a wide variation fromday to day in the same locality because of the winds and changes intemperature. ~Processes affecting the composition of the air. ~ The most important ofthese processes are the following. 1. _Respiration. _ In the process of respiration some of the oxygen inthe inhaled air is absorbed by the blood and carried to all parts of thebody, where it combines with the carbon of the worn-out tissues. Theproducts of oxidation are carried back to the lungs and exhaled in theform of carbon dioxide. The amount exhaled by an adult averages about 20l. Per hour. Hence in a poorly ventilated room occupied by a number ofpeople the amount of carbon dioxide rapidly increases. While this gas isnot poisonous unless present in large amounts, nevertheless aircontaining more than 15 parts in 10, 000 is not fit for respiration. 2. _Combustion. _ All of the ordinary forms of fuel contain largepercentages of carbon. On burning, this carbon combines with oxygen inthe air, forming carbon dioxide. Combustion and respiration, therefore, tend to diminish the amount of oxygen in the air and to increase theamount of carbon dioxide. 3. _Action of plants. _ Plants have the power, when in the sunlight, ofabsorbing carbon dioxide from the air, retaining the carbon andreturning at least a portion of the oxygen to the air. It will beobserved that these changes are just the opposite of those brought aboutby the processes of respiration and combustion. ~Poisonous effect of exhaled air. ~ The differences in the percentages of oxygen, carbon dioxide, and moisture present in inhaled air and exhaled air are shown in the following analyses. INHALED AIR EXHALED AIR Oxygen 21. 00% 16. 00% Carbon dioxide 0. 04 4. 38 Moisture variable saturated The foul odor of respired air is due to the presence of a certain amount of organic matter. It is possible that this organic matter rather than the carbon dioxide is responsible for the injurious effects which follow the respiration of impure air. The extent of such organic impurities present may be judged, however, by the amount of carbon dioxide present, since the two are exhaled together. ~The cycle of carbon in nature. ~ Under the influence of sunlight, the carbon dioxide absorbed from the air by plants reacts with water and small amounts of other substances absorbed from the soil to form complex compounds of carbon which constitute the essential part of the plant tissue. This reaction is attended by the evolution of oxygen, which is restored to the air. The compounds resulting from these changes are much richer in their energy content than are the substances from which they are formed; hence a certain amount of energy must have been absorbed in their formation. The source of this energy is the sun's rays. If the plant is burned, the changes which took place in the formation of the compounds present are largely reversed. The carbon and hydrogen present combine with oxygen taken from the air to form carbon dioxide and water, while the energy absorbed from the sun's rays is liberated in the form of energy of heat. If, on the other hand, the plant is used as food, the compounds present are used in building up the tissues of the body. When this tissue breaks down, the changes which it undergoes are very similar to those which take place when the plant is burned. The carbon and hydrogen combine with the inhaled oxygen to form carbon dioxide and water, which are exhaled. The energy possessed by the complex substances is liberated partly in the form of energy of heat, which maintains the heat of the body, and partly in the various forms of muscular energy. The carbon originally absorbed from the air by the plant in the form of carbon dioxide is thus restored to the air and is ready to repeat the cycle of changes. ~The composition of the air is constant. ~ Notwithstanding the changesconstantly taking place which tend to alter the composition of the air, the results of a great many analyses of air collected in the open fieldsshow that the percentages of oxygen and nitrogen as well as of carbondioxide are very nearly constant. Indeed, so constant are thepercentages of oxygen and nitrogen that the question has arisen, whetherthese two elements are not combined in the air, forming a definitechemical compound. That the two are not combined but are simply mixedtogether can be shown in a number of ways, among which are thefollowing. 1. When air dissolves in water it has been found that the ratio ofoxygen to nitrogen in the dissolved air is no longer 21: 78, but morenearly 35: 65. If it were a chemical compound, the ratio of oxygen tonitrogen would not be changed by solution in water. 2. A chemical compound in the form of a liquid has a definite boilingpoint. Water, for example, boils at 100°. Moreover the steam which isthus formed has the same composition as the water. The boiling point ofliquid air, on the other hand, gradually rises as the liquid boils, thenitrogen escaping first followed by the oxygen. If the two werecombined, they would pass off together in the ratio in which they arefound in the air. ~Why the air has a constant composition. ~ If air is a mixture and changesare constantly taking place which tend to modify its composition, how, then, do we account for the constancy of composition which the analysesreveal? This is explained by several facts. (1) The changes which arecaused by the processes of combustion and respiration, on the one hand, and the action of plants, on the other, tend to equalize each other. (2)The winds keep the air in constant motion and so prevent local changes. (3) The volume of the air is so vast and the changes which occur are sosmall compared with the total amount of air that they cannot be readilydetected. (4) Finally it must be noted that only air collected in theopen fields shows this constancy in composition. The air in a poorlyventilated room occupied by a number of people rapidly changes incomposition. ~The properties of the air. ~ Inasmuch as air is composed principally of amixture of oxygen and nitrogen, which elements have already beendiscussed, its properties may be inferred largely from those of the twogases. One liter weighs 1. 2923 g. It is thus 14. 38 times as heavy ashydrogen. At the sea level it exerts an average pressure sufficient tosustain a column of mercury 760 mm. In height. This is taken as thestandard pressure in determining the volumes of gases as well as theboiling points of liquids. Water may be made to boil at any temperaturebetween 0° and considerably above 100° by simply varying the pressure. It is only when the pressure upon it is equal to the normal pressure ofthe atmosphere at the sea level, as indicated by a barometric reading of760 mm. , that it boils at 100°. ~Preparation of liquid air. ~ Attention has been called to the fact thatboth oxygen and nitrogen can be obtained in the liquid state by stronglycooling the gases and applying great pressure to them. Since air islargely a mixture of these two gases, it can be liquefied by the samemethods. The methods for liquefying air have been simplified greatly in that the low temperature required is obtained by allowing a portion of the compressed air to expand. The expansion of a gas is always attended by the absorption of heat. In liquefying air the apparatus is so constructed that the heat absorbed is withdrawn from air already under great pressure. This process is continued until the temperature is lowered to the point of liquefaction. [Illustration: Fig. 29] ~The Dewar bulb. ~ It is not possible to preserve air in the liquid statein a closed vessel, on account of the enormous pressure exerted by it inits tendency to pass into the gaseous state. It may however be preservedfor some hours or even days before it will completely evaporate, bysimply placing it in an open vessel surrounded by a nonconductingmaterial. The most efficient vessel for this purpose is the _Dewar bulb_shown in Fig. 29. The air is withdrawn from the space between the twowalls, thus making it nonconducting. ~Properties and uses of liquid air. ~ When first prepared, liquid air iscloudy because of the presence of particles of solid carbon dioxide. These may be filtered off, leaving a liquid of slightly bluish color. Itbegins to boil at about -190°, the nitrogen passing off first, graduallyfollowed by the oxygen, the last portions being nearly pure oxygen. To acertain extent oxygen is now prepared in this way for commercialpurposes. The extremely low temperature of liquid air may be inferred from thefact that mercury when cooled by it is frozen to a mass so hard that itmay be used for driving nails. Liquid air is used in the preparation of oxygen and as a cooling agentin the study of the properties of matter at low temperatures. It hasthus been found that elements at extremely low temperatures largely losetheir chemical activity. EXERCISES 1. When oxygen and nitrogen are mixed in the proportion in which theyexist in the atmosphere, heat is neither evolved nor absorbed by theprocess. What important point does this suggest? 2. What essential constituent of the air is found in larger amount inmanufacturing districts than in the open country? 3. Can you suggest any reason why the growth of clover in a fieldimproves the soil? 4. Why are the inner walls of a Dewar bulb sometimes coated with a filmof silver? 5. To what is the blue color of liquid air due? Does this color increasein intensity on standing? 6. When ice is placed in a vessel containing liquid air, the latterboils violently. Explain. 7. Taking the volumes of the oxygen and nitrogen in 100 volumes of airas 21 and 78 respectively, calculate the percentages of these elementspresent by weight. 8. Would combustion be more intense in liquid air than in the gaseoussubstance? 9. A tube containing calcium chloride was found to weigh 30. 1293 g. Avolume of air which weighed 15. 2134 g. Was passed through, after whichthe weight of the tube was found to be 30. 3405 g. What was thepercentage amount of moisture present in the air? 10. 10 l. Of air measured at 20° and 740 mm. Passed through lime watercaused the precipitation of 0. 0102 g. Of CaCO_{3}. Find the number ofvolumes of carbon dioxide in 10, 000 volumes of the air. CHAPTER IX SOLUTIONS ~Definitions. ~ When a substance disappears in a liquid in such a way as tothoroughly mix with it and to be lost to sight as an individual body, the resulting liquid is called a _solution_. The liquid in which thesubstance dissolves is called the _solvent_, while the dissolvedsubstance is called the _solute_. ~Classes of solutions. ~ Matter in any one of its physical states maydissolve in a liquid, so that we may have solutions of gases, ofliquids, and of solids. Solutions of liquids in liquids are not oftenmentioned in the following pages, but the other two classes will becomevery familiar in the course of our study, and deserve special attention. SOLUTION OF GASES IN LIQUIDS [Illustration: Fig. 30] It has already been stated that oxygen, hydrogen, and nitrogen areslightly soluble in water. Accurate study has led to the conclusion thatall gases are soluble to some extent not only in water but in many otherliquids. The amount of a gas which will dissolve in a liquid dependsupon a number of conditions, and these can best be understood bysupposing a vessel B (Fig. 30), to be filled with the gas and invertedover the liquid. Under these circumstances the gas cannot escape orbecome mixed with another gas. ~Circumstances affecting the solubility of gases. ~ A number ofcircumstances affect the solubility of a gas in a liquid. 1. _Nature of the gas. _ Other conditions being equal, each gas has itsown peculiar solubility, just as it has its own special taste or odor. The solubility of gases varies between wide limits, as will be seen fromthe following table, but as a rule a given volume of a liquid will notdissolve more than two or three times its own volume of a gas. _Solubility of Gases in Water_ 1 l. Of water at 760 mm. Pressure and at 0° will dissolve: Ammonia 1148. 00 l. Hydrochloric acid 503. 00 Sulphur dioxide 79. 79 Carbon dioxide 1. 80 Oxygen 41. 14 cc. Hydrogen 21. 15 Nitrogen 20. 03 In the case of very soluble gases, such as the first three in the table, it is probable that chemical combination between the liquid and the gastakes place. 2. _Nature of the liquid. _ The character of the liquid has muchinfluence upon the solubility of a gas. Water, alcohol, and ether haveeach its own peculiar solvent power. From the solubility of a gas inwater, no prediction can be made as to its solubility in other liquids. 3. _Influence of pressure. _ It has been found that the weight of gaswhich dissolves in a given case is proportional to the pressure exertedupon the gas. If the pressure is doubled, the weight of gas going intosolution is doubled; if the pressure is diminished to one half of itsoriginal value, half of the dissolved gas will escape. Under highpressure, large quantities of gas can be dissolved in a liquid, and whenthe pressure is removed the gas escapes, causing the liquid to foam or_effervesce_. 4. _Influence of temperature. _ In general, the lower the temperature ofthe liquid, the larger the quantity of gas which it can dissolve. 1000volumes of water at 0° will dissolve 41. 14 volumes of oxygen; at 50°, 18. 37 volumes; at 100° none at all. While most gases can be expelledfrom a liquid by boiling the solution, some cannot. For example, it isnot possible to expel hydrochloric acid gas completely from its solutionby boiling. SOLUTION OF SOLIDS IN LIQUIDS This is the most familiar class of solutions, since in the laboratorysubstances are much more frequently used in the form of solutions thanin the solid state. ~Circumstances affecting the solubility of a solid. ~ The solubility of asolid in a liquid depends upon several factors. 1. _Nature of the solid. _ Other conditions being the same, solids varygreatly in their solubility in liquids. This is illustrated in thefollowing table: _Table of Solubility of Solids at 18°_ 100 cc. Of water will dissolve: Calcium chloride 71. 0 g. Sodium chloride 35. 9 Potassium nitrate 29. 1 Copper sulphate 21. 4 Calcium sulphate 0. 207 No solids are absolutely insoluble, but the amount dissolved may be sosmall as to be of no significance for most purposes. Thus bariumsulphate, one of the most insoluble of common substances, dissolves inwater to the extent of 1 part in 400, 000. 2. _Nature of the solvent. _ Liquids vary much in their power to dissolvesolids. Some are said to be good solvents, since they dissolve a greatvariety of substances and considerable quantities of them. Others havesmall solvent power, dissolving few substances, and those to a slightextent only. Broadly speaking, water is the most general solvent, andalcohol is perhaps second in solvent power. 3. _Temperature. _ The weight of a solid which a given liquid candissolve varies with the temperature. Usually it increases rapidly asthe temperature rises, so that the boiling liquid dissolves severaltimes the weight which the cold liquid will dissolve. In some instances, as in the case of common salt dissolved in water, the temperature haslittle influence upon the solubility, and a few solids are more solublein cold water than in hot. The following examples will serve asillustrations: _Table of Solubility at 0° and at 100°_ 100 cc. Of water will dissolve: At 0° At 100° Calcium chloride 49. 6 g. 155. 0 g. Sodium chloride 35. 7 39. 8 Potassium nitrate 13. 3 247. 0 Copper sulphate 15. 5 73. 5 Calcium sulphate 0. 205 0. 217 Calcium hydroxide 0. 173 0. 079 ~Saturated solutions. ~ A liquid will not dissolve an unlimited quantity ofa solid. On adding the solid to the liquid in small portions at a time, it will be found that a point is reached at which the liquid will notdissolve more of the solid at that temperature. The solid and thesolution remain in contact with each other unchanged. This condition maybe described by saying that they are in equilibrium with each other. Asolution is said to be _saturated_ when it remains unchanged inconcentration in contact with some of the solid. The weight of the solidwhich will completely saturate a definite volume of a liquid at a giventemperature is called the _solubility_ of the substance at thattemperature. ~Supersaturated solutions. ~ When a solution, saturated at a giventemperature, is allowed to cool it sometimes happens that no solidcrystallizes out. This is very likely to occur when the vessel used isperfectly smooth and the solution is not disturbed in any way. Such asolution is said to be _supersaturated_. That this condition is unstablecan be shown by adding a crystal of the solid to the solution. All ofthe solid in excess of the quantity required to saturate the solution atthis temperature will at once crystallize out, leaving the solutionsaturated. Supersaturation may also be overcome in many cases byvigorously shaking or stirring the solution. ~General physical properties of solutions. ~ A few general statements maybe made in reference to the physical properties of solutions. 1. _Distribution of the solid in the liquid. _ A solid, when dissolved, tends to distribute itself uniformly through the liquid, so that everypart of the solution has the same concentration. The process goes onvery slowly unless hastened by stirring or shaking the solution. Thus, if a few crystals of a highly colored substance such as copper sulphateare placed in the bottom of a tall vessel full of water, it will takeweeks for the solution to become uniformly colored. 2. _Boiling points of solutions. _ The boiling point of a liquid israised by the presence of a substance dissolved in it. In general theextent to which the boiling point of a solvent is raised by a givensubstance is proportional to the concentration of the solution, thatis, to the weight of the substance dissolved in a definite weight of thesolvent. 3. _Freezing points of solutions. _ A solution freezes at a lowertemperature than the pure solvent. The lowering of the freezing pointobeys the same law which holds for the raising of the boiling point: theextent of lowering is proportional to the weight of dissolved substance, that is, to the concentration of the solution. ~Electrolysis of solutions. ~ Pure water does not appreciably conduct theelectric current. If, however, certain substances such as common saltare dissolved in the water, the resulting solutions are found to beconductors of electricity. Such solutions are called _electrolytes_. When the current passes through an electrolyte some chemical changealways takes place. This change is called _electrolysis_. [Illustration: Fig. 31] The general method used in the electrolysis of a solution is illustratedin Fig. 31. The vessel D contains the electrolyte. Two plates or rods, A and B, made of suitable material, are connected with the wiresfrom a battery (or dynamo) and dipped into the electrolyte, as shown inthe figure. These plates or rods are called _electrodes_. The electrodeconnected with the zinc plate of the battery is the negative electrodeor _cathode_, while that connected with the carbon plate is the positiveelectrode or _anode_. ~Theory of electrolytic dissociation. ~ The facts which have just beendescribed in connection with solutions, together with many others, haveled chemists to adopt a theory of solutions called _the theory ofelectrolytic dissociation_. The main assumptions in this theory are thefollowing. 1. _Formation of ions. _ Many compounds when dissolved in water undergoan important change. A portion of their molecules fall apart, or_dissociate_, into two or more parts, called _ions_. Thus sodium nitrate(NaNO_{3}) dissociates into the ions Na and NO_{3}; sodium chloride, into the ions Na and Cl. These ions are free to move about in thesolution independently of each other like independent molecules, and forthis reason were given the name ion, which signifies a wanderer. 2. _The electrical charge of ions. _ Each ion carries a heavy electricalcharge, and in this respect differs from an atom or molecule. It isevident that the sodium in the form of an ion must differ in someimportant way from ordinary sodium, for sodium ions, formed from sodiumnitrate, give no visible evidence of their presence in water, whereasmetallic sodium at once decomposes the water. The electrical charge, therefore, greatly modifies the usual chemical properties of theelement. 3. _The positive charges equal the negative charges. _ The ions formed bythe dissociation of any molecule are of two kinds. One kind is chargedwith positive electricity and the other with negative electricity;moreover the sum of all the positive charges is always equal to the sumof all the negative charges. The solution as a whole is thereforeelectrically neutral. If we represent dissociation by the usual chemicalequations, with the electrical charges indicated by + and - signsfollowing the symbols, the dissociation of sodium chloride molecules isrepresented thus: NaCl --> Na^{+}, Cl^{-}. The positive charge on each sodium ion exactly equals the negativecharge on each chlorine ion. Sodium sulphate dissociates, as shown inthe equation Na_{2}SO_{4} --> 2Na^{+}, SO_{4}^{--}. Here the positive charge on the two sodium ions equals the doublenegative charge on the SO_{4} ion. 4. _Not all compounds dissociate. _ Only those compounds dissociate whosesolutions form electrolytes. Thus salt dissociates when dissolved inwater, the resulting solution being an electrolyte. Sugar, on the otherhand, does not dissociate and its solution is not a conductor of theelectric current. 5. _Extent of dissociation differs in different liquids. _ Whilecompounds most readily undergo dissociation in water, yet dissociationoften occurs to a limited extent when solution takes place in liquidsother than water. In the discussion of solutions it will be understoodthat the solvent is water unless otherwise noted. ~The theory of electrolytic dissociation and the properties of solutions. ~In order to be of value, this theory must give a reasonable explanationof the properties of solutions. Let us now see if the theory is inharmony with certain of these properties. ~The theory of electrolytic dissociation and the boiling and freezingpoints of solutions. ~ We have seen that the boiling point of a solutionof a substance is raised in proportion to the concentration of thedissolved substance. This is but another way of saying that the changein the boiling point of the solution is proportional to the number ofmolecules of the dissolved substance present in the solution. It has been found, however, that in the case of electrolytes the boilingpoint is raised more than it should be to conform to this law. If thesolute dissociates into ions, the reason for this becomes clear. Eachion has the same effect on the boiling point as a molecule, and sincetheir number is greater than the number of molecules from which theywere formed, the effect on the boiling point is abnormally great. In a similar way, the theory furnishes an explanation of the abnormallowering of the freezing point of electrolytes. ~The theory of electrolytic dissociation and electrolysis. ~ The changestaking place during electrolysis harmonize very completely with thetheory of dissociation. This will become clear from a study of thefollowing examples. [Illustration: Fig. 32] 1. _Electrolysis of sodium chloride. _ Fig. 32 represents a vessel inwhich the electrolyte is a solution of sodium chloride (NaCl). Accordingto the dissociation theory the molecules of sodium chloride dissociateinto the ions Na^{+} and Cl^{-}. The Na^{+} ions are attracted to thecathode owing to its large negative charge. On coming into contact withthe cathode, the Na^{+} ions give up their positive charge and are thenordinary sodium atoms. They immediately decompose the water according tothe equation Na + H_{2}O = NaOH + H, and hydrogen is evolved about the cathode. The chlorine ions on being discharged at the anode in similar manner mayeither be given off as chlorine gas, or may attack the water, asrepresented in the equation 2Cl + H_{2}O = 2HCl + O. 2. _Electrolysis of water. _ The reason for the addition of sulphuricacid to water in the preparation of oxygen and hydrogen by electrolysiswill now be clear. Water itself is not an electrolyte to an appreciableextent; that is, it does not form enough ions to carry a current. Sulphuric acid dissolved in water is an electrolyte, and dissociatesinto the ions 2 H^{+} and SO_{4}^{--}. In the process of electrolysis ofthe solution, the hydrogen ions travel to the cathode, and on beingdischarged escape as hydrogen gas. The SO_{4} ions, when discharged atthe anode, act upon water, setting free oxygen and once more formingsulphuric acid: SO_{4} + H_{2}O = H_{2}SO_{4} + O. The sulphuric acid can again dissociate and the process repeat itself aslong as any water is left. Hence the hydrogen and oxygen set free in theelectrolysis of water really come directly from the acid but indirectlyfrom the water. 3. _Electrolysis of sodium sulphate. _ In a similar way, sodium sulphate(Na_{2}SO_{4}), when in solution, gives the ions 2 Na^{+} andSO_{4}^{--}. On being discharged, the sodium atoms decompose water aboutthe cathode, as in the case of sodium chloride, while the SO_{4} ionswhen discharged at the anode decompose the water, as represented in theequation SO_{4} + H_{2}O = H_{2}SO_{4} + O [Illustration: Fig. 33] That new substances are formed at the cathode and anode may be shown inthe following way. A U-tube, such as is represented in Fig. 33, ispartially filled with a solution of sodium sulphate, and the liquid inone arm is colored with red litmus, that in the other with blue litmus. An electrode placed in the red solution is made to serve as cathode, while one in the blue solution is made the anode. On allowing thecurrent to pass, the blue solution turns red, while the red solutionturns blue. These are exactly the changes which would take place ifsodium hydroxide and sulphuric acid were to be set free at theelectrodes, as required by the theory. ~The properties of electrolytes depend upon the ions present. ~ When asubstance capable of dissociating into ions is dissolved in water, theproperties of the solution will depend upon two factors: (1) the ionsformed from the substance; (2) the undissociated molecules. Since theions are usually more active chemically than the molecules, most of thechemical properties of an electrolyte are due to the ions rather than tothe molecules. The solutions of any two substances which give the same ion will havecertain properties in common. Thus all solutions containing the copperion (Cu^{++}) are blue, unless the color is modified by the presence ofions or molecules having some other color. EXERCISES 1. Distinguish clearly between the following terms: electrolysis, electrolyte, electrolytic dissociation, ions, solute, solvent, solution, saturated solution, and supersaturated solution. 2. Why does the water from some natural springs effervesce? 3. (a) Why does not the water of the ocean freeze? (b) Why will iceand salt produce a lower temperature than ice alone? 4. Why does shaking or stirring make a solid dissolve more rapidly in aliquid? 5. By experiment it was found that a certain volume of water wassaturated at 100° with 114 g. Of potassium nitrate. On cooling to 0° aportion of the substance crystallized. (a) How many grams of thesubstance remained in solution? (b) What was the strength of thesolution at 18°? (c) How much water had been used in the experiment? 6. (a) 10 g. Of common salt were dissolved in water and the solutionevaporated to dryness; what weight of solid was left? (b) 10 g. Ofzinc were dissolved in hydrochloric acid and the solution evaporated todryness; what weight of solid was left? 7. Account for the fact that sugar sometimes deposits from molasses, even when no evaporation has taken place. 8. (a) From the standpoint of the theory of electrolytic dissociation, write the simple equation for a dilute solution of copper sulphate(CuSO_{4}); this solution is blue. (b) In the same manner, write onefor sodium sulphate; this solution is colorless. (c) How would youaccount for the color of the copper sulphate solution? 9. (a) As in the preceding exercise, write a simple equation for adilute solution of copper chloride (CuCl_{2}); this solution is blue. (b) In the same manner, write one for sodium chloride; this solutionis colorless. To what is the blue color due? 10. What component is present in concentrated sulphuric acid that isalmost wanting in very dilute sulphuric acid? 11. Why will vegetables cook faster when boiled in strong salt waterthan when boiled in pure water? 12. How do you explain the foaming of soda water? CHAPTER X ACIDS, BASES, AND SALTS; NEUTRALIZATION ~Acids, bases, and salts. ~ The three classes of compounds knownrespectively as acids, bases, and salts include the great majority ofthe compounds with which we shall have to deal. It is important, therefore, for us to consider each of these classes in a systematic way. The individual members belonging to each class will be discussed indetail in the appropriate places, but a few representatives of eachclass will be described in this chapter with special reference to thecommon properties in accordance with which they are classified. ~The familiar acids. ~ _Hydrochloric acid_ is a gas composed of hydrogen andchlorine, and has the formula HCl. The substance is very soluble inwater, and it is this solution which is usually called hydrochloricacid. _Nitric acid_ is a liquid composed of hydrogen, nitrogen, andoxygen, having the formula HNO_{3}. As sold commercially it is mixedwith about 32% of water. _Sulphuric acid_, whose composition isrepresented by the formula H_{2}SO_{4}, is an oily liquid nearly twiceas heavy as water, and is commonly called _oil of vitriol_. ~Characteristics of acids. ~ (1) All acids contain hydrogen. (2) Whendissolved in water the molecules of the acid dissociate into two kindsof ions. One of these is always hydrogen and is the cation (+), whilethe other consists of the remainder of the molecule and is the anion(-). (3) The solution tastes sour. (4) It has the power to change thecolor of certain substances called _indicators_. Thus blue litmus ischanged to red, and yellow methyl orange is changed to red. Since allacids produce hydrogen cations, while the anions of each are different, the properties which all acids have in common when in solution, such astaste and action on indicators, must be attributed to the hydrogen ions. DEFINITION: _An acid is a substance which produces hydrogen ions whendissolved in water or other dissociating liquids. _ ~Undissociated acids. ~ When acids are perfectly free from water, or aredissolved in liquids like benzene which do not have the power ofdissociating them into ions, they should have no real acid properties. This is found to be the case. Under these circumstances they do notaffect the color of indicators or have any of the propertiescharacteristic of acids. The familiar bases. The bases most used in the laboratory are sodiumhydroxide (NaOH), potassium hydroxide (KOH), and calcium hydroxide(Ca(OH)_{2}). These are white solids, soluble in water, the lattersparingly so. Some bases are very difficultly soluble in water. The verysoluble ones with most pronounced basic properties are sometimes calledthe _alkalis_. ~Characteristics of bases. ~ (1) All bases contain hydrogen and oxygen. (2)When dissolved in water the molecules of the base dissociate into twokinds of ions. One of these is always composed of oxygen and hydrogenand is the anion. It has the formula OH and is called the _hydroxylion_. The remainder of the molecule, which usually consists of a singleatom, is the cation. (3) The solution of a base has a soapy feel and abrackish taste. (4) It reverses the color change produced in indicatorsby acids, turning red litmus blue, and red methyl orange yellow. Sinceall bases produce hydroxyl anions, while the cations of each aredifferent, the properties which all bases have in common when insolution must be due to the hydroxyl ions. DEFINITION: _A base is a substance which produces hydroxyl ions whendissolved in water or other dissociating liquids. _ ~Undissociated bases. ~ Bases, in the absence of water or when dissolved inliquids which do not dissociate them, should have none of the propertiescharacteristic of this class of substances. This has been found to bethe case. For example, they have no effect upon indicators under thesecircumstances. ~Neutralization. ~ When an acid and a base are brought together in solutionin proper proportion, the characteristic properties of each disappear. The solution tastes neither sour nor brackish; it has no effect uponindicators. There can therefore be neither hydrogen nor hydroxyl ionspresent in the solution. A study of reactions of this kind has shownthat the hydrogen ions of the acid combine with the hydroxyl ions of thebase to form molecules of water, water being a substance which is notappreciably dissociated into ions. This action of an acid on a base iscalled _neutralization_. The following equations express theneutralization of the three acids by three bases, water being formed ineach case. Na^{+}, OH^{-} + H^{+}, Cl^{-} = Na^{+}, Cl^{-} + H_{2}O. K^{+}, OH^{-} + H^{+}, NO_{3}^{-} = K^{+}, NO_{3}^{-} + H_{2}O. Ca^{++}, (OH)_{2}^{--} + H_{2}^{++}, SO_{4}^{--} = Ca^{++}, SO_{4}^{--} + 2H_{2}O. DEFINITION: _Neutralization consists in the union of the hydrogen ion ofan acid with the hydroxyl ion of a base to form water. _ ~Salts. ~ It will be noticed that in neutralization the anion of the acidand the cation of the base are not changed. If, however, the water isexpelled by evaporation, these two ions slowly unite, and when the waterbecomes saturated with the substance so produced, it separates in theform of a solid called a _salt_. DEFINITION: _A salt is a substance formed by the union of the anion ofan acid with the cation of a base. _ ~Characteristics of salts. ~ (1) From the definition of a salt it will beseen that there is no element or group of elements which characterizesalts. (2) Salts as a class have no peculiar taste. (3) In the absenceof all other substances they are without action on indicators. (4) Whendissolved in water they form two kinds of ions. ~Heat of neutralization. ~ If neutralization is due to the union of hydrogen ions with hydroxyl ions, and nothing more, it follows that when a given weight of water is formed in neutralization, the heat set free should always be the same, no matter from what acid and base the two kinds of ions have been supplied. Careful experiments have shown that this is the case, provided no other reactions take place at the same time. When 18g. Of water are formed in neutralization, 13, 700 cal. Of heat are set free. This is represented in the equations Na^{+}, OH^{-} + H^{+}, Cl^{-} = Na^{+}, Cl^{-} + H_{2}O + 13, 700 cal. K^{+}, OH^{-} + H^{+}, NO_{3}^{-} = K^{+}, NO_{3}^{-} + H_{2}O + 13, 700 cal. Ca^{++}, (OH)_{2}^{--} + H_{2}^{++}, SO_{4}^{--} = Ca^{++}, SO_{4}^{--} + 2H_{2}O + 2 × 13, 700 cal. ~Neutralization a quantitative act. ~ Since neutralization is a definitechemical act, each acid will require a perfectly definite weight of eachbase for its neutralization. For example, a given weight of sulphuricacid will always require a definite weight of sodium hydroxide, inaccordance with the equation H_{2}, SO_{4} + 2Na, OH = Na_{2}, SO_{4} + 2H_{2}O. ~Determination of the ratio in neutralization. ~ The quantities of acid and base required in neutralization may be determined in the following way. Dilute solutions of the two substances are prepared, the sulphuric acid being placed in one of the burettes (Fig. 34) and the sodium hydroxide in the other. The levels of the two liquids are then brought to the zero marks of the burettes by means of the stopcocks. A measured volume of the acid is drawn off into a beaker, a few drops of litmus solution added, and the sodium hydroxide is run in drop by drop until the red litmus just turns blue. The volume of the sodium hydroxide consumed is then noted. If the concentrations of the two solutions are known, it is easy to calculate what weight of sodium hydroxide is required to neutralize a given weight of sulphuric acid. By evaporating the neutralized solution to dryness, the weight of the sodium sulphate formed can be determined directly. Experiment shows that the weights are always in accordance with the equation in the preceding paragraph. [Illustration: Fig. 34] ~Extent of dissociation. ~ The question will naturally arise, When an acid, base, or salt dissolves in water, do all the molecules dissociate intoions, or only a part of them? The experiments by which this question canbe answered cannot be described here. It has been found, however, thatonly a fraction of the molecules dissociate. The percentage which willdissociate in a given case depends upon several conditions, the chief ofwhich are: (1) The concentration of the solution. In concentratedsolutions only a very small percentage of dissociation occurs. As thesolution is diluted the percentage increases, and in very dilutesolutions it may be very large, though it is never complete in anyordinary solution. (2) The nature of the dissolved compound. At equalconcentrations substances differ much among themselves in the percentageof dissociation. The great majority of salts are about equallydissociated. Acids and bases, on the contrary, show great differences. Some are freely dissociated, while others are dissociated to but aslight extent. ~Strength of acids and bases. ~ Since acid and basic properties are due tohydrogen and hydroxyl ions respectively, the acid or base which willproduce the greatest percentage of these ions at a given concentrationmust be regarded as the strongest representative of its class. The acidsand bases described in the foregoing paragraphs are all quite strong. In10% solutions they are dissociated to about 50%, and this is alsoapproximately the extent to which most salts are dissociated at thissame concentration. ~Partial neutralization. ~ 1. _Basic salts. _ The chemical action between an acid and a base is not always as complete as has been represented in the foregoing paragraphs. For example, if the base magnesium hydroxide (Mg(OH)_{2}) and hydrochloric acid (HCl) are brought together in the ratio of an equal number of molecules of each, there will be only half enough hydrogen ions for the hydroxyl ions present. Mg, (OH)_{2} + H, Cl = Mg, OH, Cl + H_{2}O. Magnesium, hydroxyl, and chlorine ions are left at the close of the reaction, and under the proper conditions unite to form molecules of the compound Mg(OH)Cl. This compound, when dissolved, can form hydroxyl ions and therefore possesses basic properties; it can also form the ions of a salt (Mg and Cl), and has properties characteristic of salts. Substances of this kind are called _basic salts. _ DEFINITION: _A basic salt is a substance which can give the ions both of a base and of a salt when dissolved in water. _ 2. _Acid salts. _ In a similar way, when sulphuric acid and sodium hydroxide are brought together in the ratio of equal numbers of the molecules of each, it is possible to have a reaction expressed by the equation Na, OH + H_{2}, SO_{4} = Na, H, SO_{4} + H_{2}O. The ions remaining after all the hydroxyl ions have been used up are those of an acid (H) and those of a salt (Na and SO_{4}). These unite to form the substance NaHSO_{4}, and as the solution becomes saturated with this substance through evaporation, it separates in the form of crystals. In solution this substance can give hydrogen ions, and therefore possesses acid properties; it can also give the ions characteristic of a salt. It is therefore called an _acid salt_. DEFINITION: _An acid salt is one which can give the ions of an acid and of a salt when in solution. _ 3. _Normal salts. _ Salts which are the products of complete neutralization, such as Na_{2}SO_{4}, and which in solution can give neither hydrogen nor hydroxyl ions, but only the ions of a salt, are called _normal salts_ to distinguish them from acid and basic salts. ~Methods of expressing reactions between compounds in solution. ~ Chemicalequations representing reactions between substances in solution mayrepresent the details of the reaction, or they may simply indicate thefinal products formed. In the latter case the formation of ions is notindicated. Thus, if we wish to call attention to the details of thereaction between sodium hydroxide and hydrochloric acid in solution, theequation is written as follows: Na^{+}, OH^{-} + H^{+}, Cl^{-} = Na^{+}, Cl^{-} + H_{2}O. On the other hand, if we wish simply to represent the final productsformed, the following is used. NaOH + HCl = NaCl + H_{2}O. Both of these methods will therefore be used: ~Radicals. ~ It has been emphasized that the hydroxyl group (OH) alwaysforms the anion of a base, while the group NO_{3} forms the anion ofnitric acid and sodium nitrate; the group SO_{4}, the anion of sulphuricacid and calcium sulphate. A group of elements which in this wayconstitutes a part of a molecule, acting as a unit in a chemical change, or forming ions in solution, is called a _radical_. Some of theseradicals have been given special names, the names signifying theelements present in the radical. Thus we have the hydroxyl radical (OH)and the nitrate radical (NO_{3}). DEFINITION: _A radical is a group of elements forming part of amolecule, and acting as a unit in chemical reactions. _ ~Names of acids, bases, and salts. ~ Since acids, bases, and salts are sointimately related to each other, it is very advantageous to give namesto the three classes in accordance with some fixed system. The systemuniversally adopted is as follows: ~Naming of bases. ~ All bases are called _hydroxides_. They aredistinguished from each other by prefixing the name of the element whichis in combination with the hydroxyl group. Examples: sodium hydroxide(NaOH); calcium hydroxide (Ca(OH)_{2}); copper hydroxide (Cu(OH)_{2}). ~Naming of acids. ~ The method of naming acids depends upon whether theacid consists of two elements or three. 1. _Binary acids. _ Acids containing only one element in addition tohydrogen are called _binary acids_. They are given names consisting ofthe prefix _hydro-_, the name of the second element present, and thetermination _-ic_. Examples: hydrochloric acid (HCl); hydrosulphuricacid (H_{2}S). 2. _Ternary acids. _ In addition to the two elements present in binaryacids, the great majority of acids also contain oxygen. They thereforeconsist of three elements and are called _ternary acids_. It usuallyhappens that the same three elements can unite in different proportionsto make several different acids. The most familiar one of these is givena name ending in the suffix _-ic_, while the one with less oxygen isgiven a similar name, but ending in the suffix _-ous_. Examples: nitricacid (HNO_{3}); nitrous acid (HNO_{2}). In cases where more than twoacids are known, use is made of prefixes in addition to the two suffixes_-ic_ and _-ous_. Thus the prefix _per-_ signifies an acid still richerin oxygen; the prefix _hypo-_ signifies one with less oxygen. ~Naming of salts. ~ A salt derived from a binary acid is given a nameconsisting of the names of the two elements composing it, with thetermination _-ide_. Example: sodium chloride (NaCl). All other binarycompounds are named in the same way. A salt of a ternary acid is named in accordance with the acid from whichit is derived. A ternary acid with the termination _-ic_ gives a saltwith the name ending in _-ate_, while an acid with termination _-ous_gives a salt with the name ending in _-ite_. The following table willmake the application of these principles clear: ACIDS SYMBOL SALTS SYMBOL Hydrochloric HCl Sodium chloride NaCl Hypochlorous HClO Sodium hypochlorite NaClO Chlorous HClO_{2} Sodium chlorite NaClO_{2} Chloric HClO_{3} Sodium chlorate NaClO_{3} Perchloric HClO_{4} Sodium perchlorate NaClO_{4} EXERCISES 1. 25 cc. Of a solution containing 40 g. Of sodium hydroxide per literwas found to neutralize 25 cc. Of a solution of hydrochloric acid. Whatwas the strength of the acid solution? 2. After neutralizing a solution of sodium hydroxide with nitric acid, there remained after evaporation 100 g. Of sodium nitrate. How much ofeach substance had been used? 3. A solution contains 18 g. Of hydrochloric acid per 100 cc. Itrequired 25 cc. Of this solution to neutralize 30 cc. Of a solution ofsodium hydroxide. What was the strength of the sodium hydroxide solutionin parts per hundred? 4. When perfectly dry sulphuric acid is treated with perfectly drysodium hydroxide, no chemical change takes place. Explain. 5. When cold, concentrated sulphuric acid is added to zinc, no changetakes place. Recall the action of dilute sulphuric acid on the samemetal. How do you account for the difference? 6. A solution of hydrochloric acid in benzene does not conduct theelectric current. When this solution is treated with zinc, will hydrogenbe evolved? Explain. 7. (a) Write equation for preparation of hydrogen from zinc and dilutesulphuric acid. (b) Rewrite the same equation from the standpoint ofthe theory of electrolytic dissociation, (c) Subtract the commonSO_{4} ion from both members of the equation, (d) From the resultingequation, explain in what the preparation of hydrogen consists whenexamined from the standpoint of this theory. 8. In the same manner as in the preceding exercise, explain in what theaction of sodium on water to give hydrogen consists. CHAPTER XI VALENCE ~Definition of valence. ~ A study of the formulas of various binarycompounds shows that the elements differ between themselves in thenumber of atoms of other elements which they are able to hold incombination. This is illustrated in the formulas HCl, H_{2}O, H_{3}N, H_{4}C. (hydrochloric acid) (water) (ammonia) (marsh gas) It will be noticed that while one atom of chlorine combines with oneatom of hydrogen, an atom of oxygen combines with two, an atom ofnitrogen with three, one of carbon with four. The number which expressesthis combining ratio between atoms is a definite property of eachelement and is called its _valence_. DEFINITION: _The valence of an element is that property which determinesthe number of the atoms of another element which its atom can hold incombination. _ ~Valence a numerical property. ~ Valence is therefore merely a numericalrelation and does not convey any information in regard to the intensityof the affinity between atoms. Judging by the heat liberated in theirunion, oxygen has a far stronger affinity for hydrogen than doesnitrogen, but an atom of oxygen can combine with two atoms only ofhydrogen, while an atom of nitrogen can combine with three. ~Measure of valence. ~ In expressing the valence of an element we mustselect some standard for comparison, just as in the measurement of anyother numerical quantity. It has been found that an atom of hydrogen isnever able to hold in combination more than one atom of any otherelement. Hydrogen is therefore taken as the standard, and other elementsare compared with it in determining their valence. A number of otherelements are like hydrogen in being able to combine with at most oneatom of other elements, and such elements are called _univalent_. Amongthese are chlorine, iodine, and sodium. Elements such as oxygen, calcium, and zinc, which can combine with two atoms of hydrogen or otherunivalent elements, are said to be _divalent_. Similarly, we have_trivalent, tetravalent, pentavalent_ elements. None have a valence ofmore than 8. ~Indirect measure of valence. ~ Many elements, especially among the metals, do not readily form compounds with hydrogen, and their valence is noteasy to determine by direct comparison with the standard element. Theseelements, however, combine with other univalent elements, such aschlorine, and their valence can be determined from the compounds soformed. ~Variable valence. ~ Many elements are able to exert different valencesunder differing circumstances. Thus we have the compounds Cu_{2}O andCuO, CO and CO_{2}, FeCl_{2} and FeCl_{3}. It is not always possible toassign a fixed valence to an element. Nevertheless each element tends toexert some normal valence, and the compounds in which it has a valencedifferent from this are apt to be unstable and easily changed intocompounds in which the valence of the element is normal. The valences ofthe various elements will become familiar as the elements are studied indetail. ~Valence and combining ratios. ~ When elements combine to form compounds, the ratio in which they combine will be determined by their valences. Inthose compounds which consist of two elements directly combined, theunion is between such numbers of the two atoms as have equal valences. Elements of the same valence will therefore combine atom for atom. Designating the valence of the atoms by Roman numerals placed abovetheir symbols, we have the formulas II II II III III IV IV HCl, ZnO, BN, CSi. A divalent element, on the other hand, will combine with two atoms of aunivalent element. Thus we have II II II II ZnCl_{2} and H_{2}O (the numerals above each symbol representing the sum of the valences ofthe atoms of the element present). A trivalent atom will combine withthree atoms of a univalent element, as in the compound III III H_{3}N. If a trivalent element combines with a divalent element, the union willbe between two atoms of the trivalent element and three of the divalentelement, since these numbers are the smallest which have equal valences. Thus the oxide of the trivalent metal aluminium has the formulaAl_{2}O_{3}. Finally one atom of a tetravalent element such as carbonwill combine with four atoms of a univalent element, as in the compoundCH_{4}, or with two atoms of a divalent element, as in the compoundCO_{2}. We have no knowledge as to why elements differ in their combining power, and there is no way to determine their valences save by experiment. ~Valence and the structure of compounds. ~ Compounds will be met from time to time which are apparent exceptions to the general statements just made in regard to valence. Thus, from the formula for hydrogen dioxide (H_{2}O_{2}), it might be supposed that the oxygen is univalent; yet it is certainly divalent in water (H_{2}O). That it may also be divalent in H_{2}O_{2} may be made clear as follows: The unit valence of each element may be represented graphically by a line attached to its symbol. Univalent hydrogen and divalent oxygen will then have the symbols H- and -O-. When atoms combine, each unit valence of one atom combines with a unit valence of another atom. Thus the composition of water may be expressed by the formula H-O-H, which is meant to show that each of the unit valences of oxygen is satisfied with the unit valence of a single hydrogen atom. The chemical conduct of hydrogen dioxide leads to the conclusion that the two oxygen atoms of its molecule are in direct combination with each other, and in addition each is in combination with a hydrogen atom. This may be expressed by the formula H-O-O-H. The oxygen in the compound is therefore divalent, just as it is in water. It will thus be seen that the structure of a compound must be known before the valences of the atoms making up the compound can be definitely decided upon. Such formulas as H-O-H and H-O-O-H are known as _structural formulas_, because they are intended to show what is known in regard to the arrangement of the atoms in the molecules. ~Valence and the replacing power of atoms. ~ Just as elements having thesame valence combine with each other atom for atom, so if they replaceeach other in a chemical reaction they will do so in the same ratio. This is seen in the following equations, in which a univalent hydrogenatom is replaced by a univalent sodium atom: NaOH + HCl = NaCl + H_{2}O. 2NaOH + H_{2}SO_{4} = Na_{2}SO_{4} + 2H_{2}O. Na + H_{2}O = NaOH + H. Similarly, one atom of divalent calcium will replace two atoms ofunivalent hydrogen or one of divalent zinc: Ca(OH)_{2} + 2 HCl = CaCl_{2} + 2H_{2}O. CaCl_{2} + ZnSO_{4} = CaSO_{4} + ZnCl_{2}. In like manner, one atom of a trivalent element will replace three of aunivalent element, or two atoms will replace three atoms of a divalentelement. ~Valence and its applications to formulas of salts. ~ While the true natureof valence is not understood and many questions connected with thesubject remain unanswered, yet many of the main facts are of much helpto the student. Thus the formula of a salt, differs from that of theacid from which it is derived in that the hydrogen of the acid has beenreplaced by a metal. If, then, it is known that a given metal forms anormal salt with a certain acid, the formula of the salt can at once bedetermined if the valence of the metal is known. Since sodium isunivalent, the sodium salts of the acids HCl and H_{2}SO_{4} will berespectively NaCl and Na_{2}SO_{4}. One atom of divalent zinc willreplace 2 hydrogen atoms, so that the corresponding zinc salts will beZnCl_{2} and ZnSO_{4}. The formula for aluminium sulphate is somewhat more difficult todetermine. Aluminium is trivalent, and the simplest ratio in which thealuminium atom can replace the hydrogen in sulphuric acid is 2 atoms ofaluminium (6 valences) to 3 molecules of sulphuric acid (6 hydrogenatoms). The formula of the sulphate will then be Al_{2}(SO_{4})_{3}. ~Valence and its application to equation writing. ~ It will be readily seenthat a knowledge of valence is also of very great assistance in writingthe equations for reactions of double decomposition. Thus, in thegeneral reaction between an acid and a base, the essential action isbetween the univalent hydrogen ion and the univalent hydroxyl ion. Thebase and the acid must always be taken in such proportions as to securean equal number of each of these ions. Thus, in the reaction betweenferric hydroxide (Fe(OH)_{3}) and sulphuric acid (H_{2}SO_{4}), it willbe necessary to take 2 molecules of the former and 3 of the latter inorder to have an equal number of the two ions, namely, 6. The equationwill then be 2Fe(OH)_{3} + 3H_{2}SO_{4} = Fe_{2}(SO_{4})_{3} + 6H_{2}O. Under certain conditions the salts Al_{2}(SO_{4})_{3} and CaCl_{2}undergo double decomposition, the two metals, aluminium and calcium, exchanging places. The simplest ratio of exchange in this case is 2atoms of aluminium (6 valences) and 3 atoms of calcium (6 valences). The reaction will therefore take place between 1 molecule ofAl_{2}(SO_{4})_{3} and 3 of CaCl_{2}, and the equation is as follows: Al_{2}(SO_{4})_{3} + 3 CaCl_{2} = 3CaSO_{4} + 2AlCl_{3}. EXERCISES 1. Sodium, calcium, and aluminium have valences of 1, 2, and 3respectively; write the formulas of their chlorides, sulphates, andphosphates (phosphoric acid = H_{3}PO_{4}), on the supposition that theyform salts having the normal composition. 2. Iron forms one series of salts in which it has a valence of 2, andanother series in which it has a valence of 3; write the formulas forthe two chlorides of iron, also for the two sulphates, on thesupposition that these have the normal composition. 3. Write the equation representing the neutralization of each of thefollowing bases by each of the acids whose formulas are given: NaOH HCl Ba(OH)_{2} H_{2}SO_{4} Al(OH)_{3} H_{3}PO_{4} 4. Silver acts as a univalent element and calcium as a divalent elementin the formation of their respective nitrates and chlorides. (a) Writethe formula for silver nitrate; for calcium chloride. (b) Whensolutions of these two salts are mixed, the two metals, silver andcalcium, exchange places; write the equation for the reaction. _5. _ Antimony acts as a trivalent element in the formation of achloride. (a) What is the formula for antimony chloride? (b) Whenhydrosulphuric acid (H_{2}S) is passed into a solution of this chloridethe hydrogen and antimony exchange places; write the equation for thereaction. 6. Lead has a valence of 2 and iron of 3 in the compounds knownrespectively as lead nitrate and ferric sulphate. (a) Write theformulas for these two compounds. (b) When their solutions are mixedthe two metals exchange places; write the equation for the reaction. CHAPTER XII COMPOUNDS OF NITROGEN ~Occurrence. ~ As has been stated in a former chapter, nitrogen constitutesa large fraction of the atmosphere. The compounds of nitrogen, however, cannot readily be obtained from this source, since at any ordinarytemperature nitrogen is able to combine directly with very few of theelements. In certain forms of combination nitrogen occurs in the soil from whichit is taken up by plants and built into complex substances composedchiefly of carbon, hydrogen, oxygen, and nitrogen. Animals feeding onthese plants assimilate the nitrogenous matter, so that this element isan essential constituent of both plants and animals. ~Decomposition of organic matter by bacteria. ~ When living matter dies andundergoes decay complicated chemical reactions take place, one result ofwhich is that the nitrogen of the organic matter is set free either asthe element nitrogen, or in the form of simple compounds, such asammonia (NH_{3}) or oxides of nitrogen. Experiment has shown that allsuch processes of decay are due to the action of different kinds ofbacteria, each particular kind effecting a different change. ~Decomposition of organic matter by heat. ~ When organic matter is stronglyheated decomposition into simpler substances takes place in much thesame way as in the case of bacterial decomposition. Coal is a complexsubstance of vegetable origin, consisting largely of carbon, but alsocontaining hydrogen, oxygen, and nitrogen. When this is heated in aclosed vessel so that air is excluded, about one seventh of the nitrogenis converted into ammonia, and this is the chief source from whichammonia and its compounds are obtained. COMPOUNDS OF NITROGEN WITH HYDROGEN ~Ammonia~ (NH_{3}). Several compounds consisting exclusively of nitrogenand hydrogen are known, but only one, ammonia, need be considered here. ~Preparation of ammonia. ~ Ammonia is prepared in the laboratory by adifferent method from the one which is used commercially. 1. _Laboratory method. _ In the laboratory ammonia is prepared fromammonium chloride, a compound having the formula NH_{4}Cl, and obtainedin the manufacture of coal gas. As will be shown later in the chapter, the group NH_{4} in this compound acts as a univalent radical and isknown as _ammonium_. When ammonium chloride is warmed with sodiumhydroxide, the ammonium and sodium change places, the reaction beingexpressed in the following equation. NH_{4}Cl + NaOH = NaCl + NH_{4}OH. The ammonium hydroxide (NH_{4}OH) so formed is unstable and breaks downinto water and ammonia. NH_{4}OH = NH_{3} + H_{2}O. Calcium hydroxide (Ca(OH)_{2}) is frequently used in place of the moreexpensive sodium hydroxide, the equations being 2NH_{4}Cl + Ca(OH)_{2} = CaCl_{2} + 2NH_{4}OH, 2NH_{4}OH = 2H_{2}O + 2NH_{3}. In the preparation, the ammonium chloride and calcium hydroxide are mixed together and placed in a flask arranged as shown in Fig. 35. The mixture is gently warmed, when ammonia is evolved as a gas and is collected by displacement of air. [Illustration: Fig. 35] 2. _Commercial method. _ Nearly all the ammonia of commerce comes fromthe gasworks. Ordinary illuminating gas is made by distilling coal, aswill be explained later, and among the products of this distillation asolution of ammonia in water is obtained. This solution, known as _gasliquor_, contains not only ammonia but other soluble substances. Most ofthese combine chemically with lime, while ammonia does not; if then limeis added to the gas liquor and the liquor is heated, the ammonia isdriven out from the mixture. It may be dissolved again in pure, coldwater, forming _aqua ammonia_, or the ammonia water of commerce. ~Preparation from hydrogen and nitrogen. ~ When electric sparks are passed for some time through a mixture of hydrogen and nitrogen, a small percentage of the two elements in the mixture is changed into ammonia. The action soon ceases, however, for the reason that ammonia is decomposed by the electric discharge. The reaction expressed in the equation N + 3H = NH_{3} can therefore go in either direction depending upon the relative quantities of the substances present. This recalls the similar change from oxygen into ozone, which soon ceases because the ozone is in turn decomposed into oxygen. ~Physical properties. ~ Under ordinary conditions ammonia is a gas whosedensity is 0. 59. It is therefore little more than half as heavy as air. It is easily condensed into a colorless liquid, and can now be purchasedin liquid form in steel cylinders. The gas is colorless and has astrong, suffocating odor. It is extremely soluble in water, 1 l. Ofwater at 0° and 760 mm. Pressure dissolving 1148 l. Of the gas. Indissolving this large volume of gas the water expands considerably, sothat the density of the solution is less than that of water, thestrongest solutions having a density of 0. 88. ~Chemical properties. ~ Ammonia will not support combustion, nor will itburn under ordinary conditions. In an atmosphere of oxygen it burns witha feeble, yellowish flame. When quite dry it is not a very activesubstance, but when moist it combines with a great many substances, particularly with acids. ~Uses. ~ It has been stated that ammonia can be condensed to a liquid bythe application of pressure. If the pressure is removed from the liquidso obtained, it rapidly passes again into the gaseous state and in sodoing absorbs a large amount of heat. Advantage is taken of this fact inthe preparation of artificial ice. Large quantities of ammonia are alsoused in the preparation of ammonium compounds. ~The manufacture of artificial ice. ~ Fig. 36 illustrates the method of preparing artificial ice. The ammonia gas is liquefied in the pipes X by means of the pump Y. The heat generated is absorbed by water flowing over the pipes. The pipes lead into a large brine tank, a cross section of which is shown in the figure. Into the brine (concentrated solution of common salt) contained in this tank are dipped the vessels A, B, C, filled with pure water. The pressure is removed from the liquid ammonia as it passes into the pipes immersed in the brine, and the heat absorbed by the rapid evaporation of the liquid lowers the temperature of the brine below zero. The water in A, B, C is thereby frozen into cakes of ice. The gaseous ammonia resulting from the evaporation of the liquid ammonia is again condensed, so that the process is continuous. [Illustration Fig. 36] ~Ammonium hydroxide~ (NH_{4}OH). The solution of ammonia in water is foundto have strong basic properties and therefore contains hydroxyl ions. Itturns red litmus blue; it has a soapy feel; it neutralizes acids, forming salts with them. It seems probable, therefore, that when ammoniadissolves in water it combines chemically with it according to theequation NH_{3} + H_{2}O = NH_{4}OH, and that it is the substance NH_{4}OH, called ammonium hydroxide, whichhas the basic properties, dissociating into the ions NH_{4} and OH. Ammonium hydroxide has never been obtained in a pure state. At everyattempt to isolate it the substance breaks up into water and ammonia, -- NH_{4}OH = NH_{3} + H_{2}O. ~The ammonium radical. ~ The radical NH_{4} plays the part of a metal inmany chemical reactions and is called ammonium. The ending _-ium_ isgiven to the name to indicate the metallic properties of the substance, since the names of the metals in general have that ending. The saltsformed by the action of the base ammonium hydroxide on acids are calledammonium salts. Thus, with hydrochloric acid, ammonium chloride isformed in accordance with the equation NH_{4}OH + HCl = NH_{4}Cl + H_{2}O. Similarly, with nitric acid, ammonium nitrate (NH_{4}NO_{3}) is formed, and with sulphuric acid, ammonium sulphate ((NH_{4})_{2}S0_{4}). It will be noticed that in the neutralization of ammonium hydroxide byacids the group NH_{4} replaces one hydrogen atom of the acid, just assodium does. The group therefore acts as a univalent metal. ~Combination of nitrogen with hydrogen by volume. ~ Under suitableconditions ammonia can be decomposed into nitrogen and hydrogen bypassing electric sparks through the gas. Accurate measurement has shownthat when ammonia is decomposed, two volumes of the gas yield one volumeof nitrogen and three volumes of hydrogen. Consequently, if the twoelements were to combine directly, one volume of nitrogen would combinewith three volumes of hydrogen to form two volumes of ammonia. Here, asin the formation of steam from hydrogen and oxygen, small whole numbersserve to indicate the relation between the volumes of combining gasesand that of the gaseous product. COMPOUNDS OF NITROGEN WITH OXYGEN AND HYDROGEN In addition to ammonium hydroxide, nitrogen forms several compounds withhydrogen and oxygen, of which nitric acid (HNO_{3}) and nitrous acid(HNO_{2}) are the most familiar. ~Nitric acid~ (HNO_{3}). Nitric acid is not found to any extent in nature, but some of its salts, especially sodium nitrate (NaNO_{3}) andpotassium nitrate (KNO_{3}) are found in large quantities. From thesesalts nitric acid can be obtained. [Illustration Fig. 37] ~Preparation of nitric acid. ~ When sodium nitrate is treated withconcentrated cold sulphuric acid, no chemical action seems to takeplace. If, however, the mixture is heated in a retort, nitric acid isgiven off as a vapor and may be easily condensed to a liquid by passingthe vapor into a tube surrounded by cold water, as shown in Fig. 37. Anexamination of the liquid left in the retort shows that it containssodium acid sulphate (NaHSO_{4}), so that the reaction may berepresented by the equation NaNO_{3} + H_{2}SO_{4} = NaHSO_{4} + HNO_{3}. If a smaller quantity of sulphuric acid is taken and the mixture is heated to a high temperature, normal sodium sulphate is formed: 2NaNO_{3} + H_{2}SO_{4} = Na_{2}SO_{4} + 2HNO_{3}. In this case, however, the higher temperature required decomposes a part of the nitric acid. ~The commercial preparation of nitric acid. ~ Fig. 38 illustrates a form of apparatus used in the preparation of nitric acid on a large scale. Sodium nitrate and sulphuric acid are heated in the iron retort A. The resulting acid vapors pass in the direction indicated by the arrows, and are condensed in the glass tubes B, which are covered with cloth kept cool by streams of water. These tubes are inclined so that the liquid resulting from the condensation of the vapors runs back into C and is drawn off into large vessels (D). [Illustration Fig. 38] ~Physical properties of nitric acid. ~ Pure nitric acid is a colorlessliquid, which boils at about 86° and has a density of 1. 56. Theconcentrated acid of commerce contains about 68% of the acid, theremainder being water. Such a mixture has a density of 1. 4. Theconcentrated acid fumes somewhat in moist air, and has a sharp chokingodor. ~Chemical properties. ~ The most important chemical properties of nitricacid are the following. 1. _Acid properties. _ As the name indicates, this substance is an acid, and has all the properties of that class of substances. It changes bluelitmus red and has a sour taste in dilute solutions. It forms hydrogenions in solution and neutralizes bases forming salts. It also acts uponthe oxides of most metals, forming a salt and water. It is one of thestrongest acids. 2. _Decomposition on heating. _ When boiled, or exposed for some time tosunlight, it suffers a partial decomposition according to the equation 2HNO_{3} = H_{2}O + 2NO_{2} + O. The substance NO_{2}, called nitrogen peroxide, is a brownish gas, whichis readily soluble in water and in nitric acid. It therefore dissolvesin the undecomposed acid, and imparts a yellowish or reddish color toit. Concentrated nitric acid highly charged with this substance iscalled _fuming nitric acid_. 3. _Oxidizing action. _ According to its formula, nitric acid contains alarge percentage of oxygen, and the reaction just mentioned shows thatthe compound is not a very stable one, easily undergoing decomposition. These properties should make it a good oxidizing agent, and we find thatthis is the case. Under ordinary circumstances, when acting as anoxidizing agent, it is decomposed according to the equation 2HNO_{3} = H_{2}O + 2NO + 3O. The oxygen is taken up by the substance oxidized, and not set free, asis indicated in the equation. Thus, if carbon is oxidized by nitricacid, the oxygen combines with carbon, forming carbon dioxide (CO_{2}): C + 2O = CO_{2}. 4. _Action on metals. _ We have seen that when an acid acts upon a metalhydrogen is set free. Accordingly, when nitric acid acts upon a metal, such as copper, we should expect the reaction to take place which isexpressed in the equation Cu + 2HNO_{3} = Cu(NO_{3})_{2} + 2H. This reaction does take place, but the hydrogen set free is immediatelyoxidized to water by another portion of the nitric acid according to theequation HNO_{3} + 3H = 2H_{2}O + NO. As these two equations are written, two atoms of hydrogen are given offin the first equation, while three are used up in the second. In orderthat the hydrogen may be equal in the two equations, we must multiplythe first by 3 and the second by 2. We shall then have 3Cu + 6HNO_{3} = 3Cu(NO_{3})_{2} + 6H, 2HNO_{3} + 6H = 4H_{2}O + 2NO. The two equations may now be combined into one by adding the quantitieson each side of the equality sign, canceling the hydrogen which is givenoff in the one reaction and used up in the other. We shall then have theequation 3Cu + 8HNO_{3} = 3Cu(NO_{3})_{2} + 2NO + 4H_{2}O. A number of other reactions may take place when nitric acid acts uponmetals, resulting in the formation of other oxides of nitrogen, freenitrogen, or even ammonia. The reaction just given is, however, theusual one. ~Importance of steps in a reaction. ~ This complete equation has the advantage of making it possible to calculate very easily the proportions in which the various substances enter into the reaction or are formed in it. It is unsatisfactory in that it does not give full information about the way in which the reaction takes place. For example, it does not suggest that hydrogen is at first formed, and subsequently transformed into water. It is always much more important to remember the steps in a chemical reaction than to remember the equation expressing the complete action; for if these steps in the reaction are understood, the complete equation is easily obtained in the manner just described. ~Salts of nitric acid, --nitrates. ~ The salts of nitric acid are callednitrates. Many of these salts will be described in the study of themetals. They are all soluble in water, and when heated to a hightemperature undergo decomposition. In a few cases a nitrate on beingheated evolves oxygen, forming a nitrite: NaNO_{3} = NaNO_{2} + O. In other cases the decomposition goes further, and the metal is left asoxide: Cu(NO_{3})_{2} = CuO + 2NO_{2} + O. ~Nitrous acid~ (HNO_{2}). It is an easy matter to obtain sodium nitrite(NaNO_{2}), as the reaction given on the previous page indicates. Instead of merely heating the nitrate, it is better to heat it togetherwith a mild reducing agent, such as lead, when the reaction takes placewhich is expressed by the equation NaNO_{3} + Pb = PbO + NaNO_{2}. When sodium nitrite is treated with an acid, such as sulphuric acid, itis decomposed and nitrous acid is set free: NaNO_{2} + H_{2}SO_{4} = NaHSO_{4} + HNO_{2}. The acid is very unstable, however, and decomposes readily into waterand nitrogen trioxide (N_{2}O_{3}): 2HNO_{2} = H_{2}O + N_{2}O_{3}. Dilute solutions of the acid, however, can be obtained. COMPOUNDS OF NITROGEN WITH OXYGEN Nitrogen combines with oxygen to form five different oxides. Theformulas and names of these are as follows: N_{2}O nitrous oxide. NO nitric oxide. NO_{2} nitrogen peroxide. N_{2}O_{3} nitrogen trioxide, or nitrous anhydride. N_{2}O_{5} nitrogen pentoxide, or nitric anhydride. These will now be briefly discussed. ~Nitrous oxide~ (_laughing gas_) (N_{2}O). Ammonium nitrate, like allnitrates, undergoes decomposition when heated; and owing to the factthat it contains no metal, but does contain both oxygen and hydrogen, the reaction is a peculiar one. It is represented by the equation NH_{4}NO_{3} = 2H_{2}O + N_{2}O. The oxide of nitrogen so formed is called nitrous oxide or laughing gas. It is a colorless gas having a slight odor. It is somewhat soluble inwater, and in solution has a slightly sweetish taste. It is easilyconverted into a liquid and can be purchased in this form. When inhaledit produces a kind of hysteria (hence the name "laughing gas"), and evenunconsciousness and insensibility to pain if taken in large amounts. Ithas long been used as an anęsthetic for minor surgical operations, suchas those of dentistry, but owing to its unpleasant after effects it isnot so much in use now as formerly. Chemically, nitrous oxide is remarkable for the fact that it is a veryenergetic oxidizing agent. Substances such as carbon, sulphur, iron, andphosphorus burn in it almost as brilliantly as in oxygen, forming oxidesand setting free nitrogen. Evidently the oxygen in nitrous oxide cannotbe held in very firm combination by the nitrogen. [Illustration Fig. 39] ~Nitric oxide~ (NO). We have seen that when nitric acid acts upon metals, such as copper, the reaction represented by the following equation takesplace: 3Cu + 8HNO_{3} = 3Cu(NO_{3})_{3} + 2NO + 4H_{2}O. Nitric oxide is most conveniently prepared in this way. The metal isplaced in the flask A (Fig. 39) and the acid added slowly through thefunnel tube B. The gas escapes through C and is collected overwater. Pure nitric oxide is a colorless gas, slightly heavier than air, and ispractically insoluble in water. It is a difficult gas to liquefy. Unlikenitrous oxide, nitric oxide does not part with its oxygen easily, andburning substances introduced into this gas are usually extinguished. Afew substances like phosphorus, which have a very strong affinity foroxygen and which are burning energetically in the air, will continue toburn in an atmosphere of nitric oxide. In this case the nitric oxideloses all of its oxygen and the nitrogen is set free as gas. ~Action of nitric oxide with oxygen. ~ When nitric oxide comes into contactwith oxygen or with the air, it at once combines with the oxygen even atordinary temperatures, forming a reddish-yellow gas of the formulaNO_{2}, which is called nitrogen peroxide. This action is not energeticenough to produce a flame, though considerable heat is set free. ~Nitrogen peroxide~ (NO_{2}). This gas, as we have just seen, is formed byallowing nitric oxide to come into contact with oxygen. It can also bemade by heating certain nitrates, such as lead nitrate: Pb(NO_{3})_{2} = PbO + 2NO_{2} + O. It is a reddish-yellow gas of unpleasant odor, which is quite poisonouswhen inhaled. It is heavier than air and is easily condensed to aliquid. It dissolves in water, but this solution is not a mere physicalsolution; the nitrogen peroxide is decomposed, forming a mixture ofnitric and nitrous acids: 2NO_{2} + H_{2}O = HNO_{2} + HNO_{3}. Nitrogen peroxide will not combine with more oxygen; it will, however, give up a part of its oxygen to burning substances, acting as anoxidizing agent: NO_{2} = NO + O. ~Acid anhydrides. ~ The oxides N_{2}O_{3} (nitrogen trioxide) andN_{2}O_{5} (nitrogen pentoxide) are rarely prepared and need not beseparately described. They bear a very interesting relation to the acidsof nitrogen. When dissolved in water they combine with the water, forming acids: N_{2}O_{3} + H_{2}O = 2HNO_{2}, N_{2}O_{5} + H_{2}O = 2HNO_{3}. On the other hand, nitrous acid very easily decomposes, yielding waterand nitrogen trioxide, and by suitable means nitric acid likewise may bedecomposed into water and nitrogen pentoxide: 2HNO_{2} = H_{2}O + N_{2}O_{3}, 2HNO_{3} = H_{2}O + N_{2}O_{5}. In view of the close relation between these oxides and the correspondingacids, they are called _anhydrides_ of the acids, N_{2}O_{3} beingnitrous anhydride and N_{2}O_{5} nitric anhydride. DEFINITION: _Any oxide which will combine with water to form an acid, orwhich together with water is formed by the decomposition of an acid, iscalled an anhydride of that acid. _ EXERCISES 1. Perfectly dry ammonia does not affect litmus paper. Explain. 2. Can ammonia be dried by passing the gas through concentratedsulphuric acid? Explain. 3. Ammonium hydroxide is a weak base, i. E. It is not highly dissociated. When it is neutralized by strong acids the heat of reaction is less thanwhen strong bases are so neutralized. Suggest some possible cause forthis. 4. Why is brine used in the manufacture of artificial ice? 5. Discuss the energy changes which take place in the manufacture ofartificial ice. 6. What weight of ammonium chloride is necessary to furnish enoughammonia to saturate 1 l. Of water at 0° and 760 mm. ? 7. What weight of sodium nitrate is necessary to prepare 100 cc. Ofcommercial nitric acid? What weight of potassium nitrate is necessary tofurnish the same weight of acid? 8. 100 l. Of nitrogen peroxide were dissolved in water and neutralizedwith sodium hydroxide. What substances were formed and how much ofeach?(1 l. Nitrogen peroxide weighs 2. 05 grams. ) 9. How many liters of nitrous oxide, measured under standard conditions, can be prepared from 10 g. Of ammonium nitrate? 10. What weight of copper is necessary to prepare 50 l. Of nitric oxideunder standard conditions? 11. (a) Calculate the percentage composition of the oxides ofnitrogen. (b) What important law does this series of substancesillustrate? 12. Write the equations representing the reactions between ammoniumhydroxide, and sulphuric acid and nitric acid respectively, inaccordance with the theory of electrolytic dissociation. 13. In the same way, write the equations representing the reactionsbetween nitric acid and each of the following bases: NaOH, KOH, NH_{4}OH, Ca(OH)_{2}. CHAPTER XIII REVERSIBLE REACTIONS AND CHEMICAL EQUILIBRIUM ~Reversible reactions. ~ The reactions so far considered have beenrepresented as continuing, when once started, until one or the othersubstance taking part in the reaction has been used up. In somereactions this is not the case. For example, we have seen that whensteam is passed over hot iron the reaction is represented by theequation 3Fe + 4H_{2}O = Fe_{3}O_{4} + 8H. On the other hand, when hydrogen is passed over hot iron oxide thereverse reaction takes place: Fe_{3}O_{4} +8H = 3Fe + 4H_{2}O. The reaction can therefore go in either direction, depending upon theconditions of the experiment. Such a reaction is called a _reversiblereaction_. It is represented by an equation with double arrows in placeof the equality sign, thus: 3Fe + 4H_{2}O Fe_{3}O_{4} + 8H. In a similar way, the equation N + 3H NH_{3} expresses the fact that under some conditions nitrogen may unite withhydrogen to form ammonia, while under other conditions ammoniadecomposes into nitrogen and hydrogen. The conversion of oxygen into ozone is also reversible and may berepresented thus: oxygen ozone. ~Chemical equilibrium. ~ Reversible reactions do not usually go on tocompletion in one direction unless the conditions under which thereaction takes place are very carefully chosen. Thus, if iron and steamare confined in a heated tube, the steam acts upon the iron, producingiron oxide and hydrogen. But these substances in turn act upon eachother to form iron and steam once more. When these two oppositereactions go on at such rates that the weight of the iron changed intoiron oxide is just balanced by the weight of the iron oxide changed intoiron, there will be no further change in the relative weights of thefour substances present in the tube. The reaction is then said to havereached an equilibrium. ~Factors which determine the point of equilibrium. ~ There are two factorswhich have a great deal of influence in determining the point at which agiven reaction will reach equilibrium. 1. _Influence of the chemical nature of the substances. _ If tworeversible reactions of the same general kind are selected, it has beenfound that the point of equilibrium is different in the two cases. Forexample, in the reactions represented by the equations 3Fe + 4H_{2}O Fe_{3}O_{4} + 8H, Zn + H_{2}O ZnO + 2H, the equilibrium will be reached when very different quantities of theiron and zinc have been changed into oxides. The individual chemicalproperties of the iron and zinc have therefore marked influence upon thepoint at which equilibrium will be reached. 2. _Influence of relative mass. _ If the tube in which the reaction 3Fe + 4H_{2}O Fe_{3}O_{4} + 8H has come to an equilibrium is opened and more steam is admitted, anadditional quantity of the iron will be changed into iron oxide. If morehydrogen is admitted, some of the oxide will be reduced to metal. Thepoint of equilibrium is therefore dependent upon the relative masses ofthe substances taking part in the reaction. When one of the substancesis a solid, however, its mass has little influence, since it is only theextent of its surface which can affect the reaction. ~Conditions under which reversible reactions are complete. ~ If, when theequilibrium between iron and steam has been reached, the tube is openedand a current of steam is passed in, the hydrogen is swept away as fastas it is formed. The opposing reaction of hydrogen upon iron oxide musttherefore cease, and the action of steam on the iron will go on untilall of the iron has been transformed into iron oxide. On the other hand, if a current of hydrogen is admitted into the tube, the steam will be swept away by the hydrogen, and all of the iron oxidewill be reduced to iron. _A reversible reaction can therefore becompleted in either direction when one of the products of the reactionis removed as fast as it is formed. _ ~Equilibrium in solution. ~ When reactions take place in solution in waterthe same general principles hold good. The matter is not so simple, however, as in the case just described, owing to the fact that many ofthe reactions in solution are due to the presence of ions. Thesubstances most commonly employed in solution are acids, bases, orsalts, and all of these undergo dissociation. Any equilibrium which maybe reached in solutions of these substances must take place between thevarious ions formed, on the one hand, and the undissociated molecules, on the other. Thus, when nitric acid is dissolved in water, equilibriumis reached in accordance with the equation H^{+} + NO_{3}^{-} HNO_{3}. ~Conditions under which reversible reactions in solution are complete. ~The equilibrium between substances in solution may be disturbed and thereaction caused to go on in one direction to completion in either ofthree ways. 1. _A gas may be formed which escapes from the solution. _ When sodiumnitrate and sulphuric acid are brought together in solution all fourions, Na^{+}, NO_{3}^{-}, H^{+}, SO_{4}^{--}, are formed. These ions arefree to rearrange themselves in various combinations. For example, theH^{+} and the NO_{3}^{-} ions will reach the equilibrium H^{+} + NO_{3}^{-} HNO_{3}. If the experiment is performed with very little water present, as is thecase in the preparation of nitric acid, the equilibrium will be reachedwhen most of the H^{+} and the NO_{3}^{-} ions have combined to formundissociated HNO_{3}. Finally, if the mixture is now heated above the boiling point of nitricacid, the acid distills away as fast as it is formed. More and moreH^{+} and NO_{3}^{-} ions will then combine, and the process willcontinue until one or the other of them has all been removed from thesolution. The substance remaining is sodium acid sulphate (NaHSO_{4}), and the reaction can therefore be expressed by the equation NaNO_{3} + H_{2}SO_{4} = NaHSO_{4} + HNO_{3}. 2. _An insoluble solid may be formed. _ When hydrochloric acid (HCl) andsilver nitrate (AgNO_{3}) are brought together in solution thefollowing ions will be present: H^{+}, Cl^{-}, Ag^{+}, NO_{3}^{-}. Theions Ag^{+} and Cl^{-} will then set up the equilibrium Ag^{+} + Cl^{-} AgCl. But silver chloride (AgCl) is almost completely insoluble in water, andas soon as a very little of it has formed the solution becomessupersaturated, and the excess of the salt precipitates. More silver andchlorine ions then unite, and this continues until practically all ofthe silver or the chlorine ions have been removed from the solution. Wethen say that the following reaction is complete: AgNO_{3} + HCl = AgCl + HNO_{3}. 3. _Two different ions may form undissociated molecules. _ In theneutralization of sodium hydroxide by hydrochloric acid the ions H^{+}and OH^{-} come to the equilibrium H^{+} + OH^{-} H_{2}O. But since water is almost entirely undissociated, equilibrium can onlybe reached when there are very few hydroxyl or hydrogen ions present. Consequently the two ions keep uniting until one or the other of them ispractically removed from the solution. When this occurs theneutralization expressed in the following equation is complete: NaOH + HCl = H_{2}O + NaCl. ~Preparation of acids. ~ The principle of reversible reactions findspractical application in the preparation of most of the common acids. Anacid is usually prepared by treating the most common of its salts withsome other acid of high boiling point. The mixture is then heated untilthe lower boiling acid desired distills out. Owing to its high boilingpoint (338°), sulphuric acid is usually employed for this purpose, mostother acids boiling below that temperature. EXERCISES 1. What would take place when solutions of silver nitrate and sodiumchloride are brought together? What other chlorides would act in thesame way? 2. Is the reaction expressed by the equation NH_{3} + H_{2}O = NH_{4}OHreversible? If so, state the conditions under which it will go in eachdirection. 3. Is the reaction expressed by the equation 2H + O = H_{2}O reversible?If so, state the conditions under which it will go in each direction. 4. Suggest a method for the preparation of hydrochloric acid. CHAPTER XIV SULPHUR AND ITS COMPOUNDS ~Occurrence. ~ The element sulphur has been known from the earliest times, since it is widely distributed in nature and occurs in large quantitiesin the uncombined form, especially in the neighborhood of volcanoes. Sicily has long been famous for its sulphur mines, and smaller depositsare found in Italy, Iceland, Mexico, and especially in Louisiana, whereit is mined extensively. In combination, sulphur occurs abundantly inthe form of sulphides and sulphates. In smaller amounts it is found in agreat variety of minerals, and it is a constituent of many animal andvegetable substances. ~Extraction of sulphur. ~ Sulphur is prepared from the native substance, the separation of crude sulphur from the rock and earthy materials withwhich it is mixed being a very simple process. The ore from the mines ismerely heated until the sulphur melts and drains away from the earthyimpurities. The crude sulphur obtained in this way is distilled in aretort-shaped vessel made of iron, the exit tube of which opens into acooling chamber of brickwork. When the sulphur vapor first enters thecooling chamber it condenses as a fine crystalline powder called_flowers of sulphur_. As the condensing chamber becomes warm, thesulphur collects as a liquid in it, and is drawn off into cylindricalmolds, the product being called _roll sulphur_ or _brimstone_. ~Physical properties. ~ Roll sulphur is a pale yellow, crystalline solid, without marked taste and with but a faint odor. It is insoluble inwater, but is freely soluble in a few liquids, notably in carbondisulphide. Roll sulphur melts at 114. 8°. Just above the melting pointit forms a rather thin, straw-colored liquid. As the temperature israised, this liquid turns darker in color and becomes thicker, until atabout 235° it is almost black and is so thick that the vessel containingit can be inverted without danger of the liquid running out. At highertemperatures it becomes thin once more, and boils at 448°, forming ayellowish vapor. On cooling the same changes take place in reverseorder. ~Varieties of sulphur. ~ Sulphur is known in two general forms, crystallineand amorphous. Each of these forms exists in definite modifications. ~Crystalline sulphur. ~ Sulphur occurs in two crystalline forms, namely, rhombic sulphur and monoclinic sulphur. 1. _Rhombic sulphur. _ When sulphur crystallizes from its solution incarbon disulphide it separates in crystals which have the same color andmelting point as roll sulphur, and are rhombic in shape. Roll sulphur ismade up of minute rhombic crystals. 2. _Monoclinic sulphur. _ When melted sulphur is allowed to cool until apart of the liquid has solidified, and the remaining liquid is thenpoured off, it is found that the solid sulphur remaining in the vesselhas assumed the form of fine needle-shaped crystals. These differ muchin appearance from the rhombic crystals obtained by crystallizingsulphur from its solution in carbon disulphide. The needle-shaped formis called _monoclinic sulphur_. The two varieties differ also in densityand in melting point, the monoclinic sulphur melting at 120°. Monoclinic and rhombic sulphur remain unchanged in contact with eachother at 96°. Above this temperature the rhombic changes intomonoclinic; at lower temperatures the monoclinic changes into rhombic. The temperature 96° is therefore called the transition point of sulphur. Heat is set free when monoclinic sulphur changes into rhombic. ~Amorphous sulphur. ~ Two varieties of amorphous sulphur can be readilyobtained. These are white sulphur and plastic sulphur. 1. _White sulphur. _ Flowers of sulphur, the preparation of which hasbeen described, consists of a mixture of rhombic crystals and amorphousparticles. When treated with carbon disulphide, the crystals dissolve, leaving the amorphous particles as a white residue. 2. _Plastic sulphur. _ When boiling sulphur is poured into cold water itassumes a gummy, doughlike form, which is quite elastic. This can beseen in a very striking manner by distilling sulphur from a small, short-necked retort, such as is represented in Fig. 40, and allowing theliquid to run directly into water. In a few days it becomes quitebrittle and passes over into ordinary rhombic sulphur. [Illustration Fig. 40] ~Chemical properties of sulphur. ~ When sulphur is heated to its kindlingtemperature in oxygen or in the air it burns with a pale blue flame, forming sulphur dioxide (SO_{2}). Small quantities of sulphur trioxide(SO_{3}) may also be formed in the combustion of sulphur. Most metalswhen heated with sulphur combine directly with it, forming metallicsulphides. In some cases the action is so energetic that the massbecomes incandescent, as has been seen in the case of iron uniting withsulphur. This property recalls the action of oxygen upon metals, and ingeneral the metals which combine readily with oxygen are apt to combinequite readily with sulphur. ~Uses of sulphur. ~ Large quantities of sulphur are used as a germicide invineyards, also in the manufacture of gunpowder, matches, vulcanizedrubber, and sulphuric acid. COMPOUNDS OF SULPHUR WITH HYDROGEN ~Hydrosulphuric acid~ (H_{2}S). This substance is a gas having thecomposition expressed by the formula H_{2}S and is commonly calledhydrogen sulphide. It is found in the vapors issuing from volcanoes, andin solution in the so-called sulphur waters of many springs. It isformed when organic matter containing sulphur undergoes decay, just asammonia is formed under similar circumstances from nitrogenous matter. ~Preparation. ~ Hydrosulphuric acid is prepared in the laboratory bytreating a sulphide with an acid. Iron sulphide (FeS) is usuallyemployed: FeS + 2HCl = FeCl_{2} + H_{2}S. A convenient apparatus is shown in Fig. 41. A few lumps of iron sulphideare placed in the bottle A, and dilute acid is added in smallquantities at a time through the funnel tube B, the gas escapingthrough the tube C. [Illustration: Fig. 41] ~Explanation of the reaction. ~ Iron sulphide is a salt of hydrosulphuric acid, and this reaction is therefore similar to the one which takes place when sulphuric acid acts upon a nitrate. In both cases a salt and an acid are brought together, and there is a tendency for the reaction to go on until a state of equilibrium is reached. This equilibrium is constantly disturbed by the escape of the gaseous acid set free, so that the reaction goes on until all of the original salt has been decomposed. The two reactions differ in that the first one is complete at ordinary temperatures, while in the case of sulphuric acid acting upon sodium nitrate, the reacting substances must be heated so as to secure a temperature at which nitric acid is a gas. ~Physical properties. ~ Hydrosulphuric acid is a colorless gas, having aweak, disagreeable taste and an exceedingly offensive odor. It is rathersparingly soluble in water at ordinary temperatures, about three volumesdissolving in one of water. In boiling water it is not soluble at all. In pure form it acts as a violent poison, and even when diluted largelywith air produces headache, dizziness, and nausea. It is a littleheavier than air, having a density of 1. 18. ~Chemical properties. ~ The most important chemical properties ofhydrosulphuric acid are the following: 1. _Acid properties. _ Hydrosulphuric acid is a weak acid. In solution inwater it turns blue litmus red and neutralizes bases, forming saltscalled _sulphides_. 2. _Action on oxygen. _ The elements composing hydrosulphuric acid haveeach a strong affinity for oxygen, and are not held together veryfirmly. Consequently the gas burns readily in oxygen or the air, according to the equation H_{2}S + 3O = H_{2}O + SO_{2}. When there is not enough oxygen for both the sulphur and the hydrogen, the latter element combines with the oxygen and the sulphur is set free: H_{2}S + O = H_{2}O + S. 3. _Reducing action. _ Owing to the ease with which hydrosulphuric aciddecomposes and the strong affinity of both sulphur and hydrogen foroxygen, the substance is a strong reducing agent, taking oxygen awayfrom many substances which contain it. 4. _Action on metals. _ Hydrosulphuric acid acts towards metals in a wayvery similar to water. Thus, when it is passed over heated iron in atube, the reaction is represented by the equation 3Fe + 4H_{2}S = Fe_{3}S_{4} + 8H. Water in the form of steam, under similar circumstances, acts accordingto the equation 3Fe + 4H_{2}O = Fe_{3}O_{4} + 8H. ~Salts of hydrosulphuric acid, --sulphides. ~ The salts of hydrosulphuricacid, called sulphides, form an important class of salts. Many of themare found abundantly in nature, and some of them are important ores. They will be frequently mentioned in connection with the metals. Most of the sulphides are insoluble in water, and some of them areinsoluble in acids. Consequently, when hydrosulphuric acid is passedinto a solution of a salt, it often happens that a sulphide isprecipitated. With copper chloride the equation is CuCl_{2} + H_{2}S = CuS + 2HCl. Because of the fact that some metals are precipitated in this way assulphides while others are not, hydrosulphuric acid is extensively usedin the separation of the metals in the laboratory. ~Explanation of the reaction. ~ When hydrosulphuric acid and copper chloride are brought together in solution, both copper and sulphur ions are present, and these will come to an equilibrium, as represented in the equation Cu^{+} + S^{-} CuS. Since copper sulphide is almost insoluble in water, as soon as a very small quantity has formed the solution becomes supersaturated, and the excess keeps precipitating until nearly all the copper or sulphur ions have been removed from the solution. With some other ions, such as iron, the sulphide formed does not saturate the solution, and no precipitate results. OXIDES OF SULPHUR Sulphur forms two well-known compounds with oxygen: sulphur dioxide(SO_{2}), sometimes called sulphurous anhydride; and sulphur trioxide(SO_{3}), frequently called sulphuric anhydride. ~Sulphur dioxide~ (SO_{2}). Sulphur dioxide occurs in nature in the gasesissuing from volcanoes, and in solution in the water of many springs. Itis likely to be found wherever sulphur compounds are undergoingoxidation. ~Preparation. ~ Three general ways may be mentioned for the preparation ofsulphur dioxide: 1. _By the combustion of sulphur. _ Sulphur dioxide is readily formed bythe combustion of sulphur in oxygen or the air: S + 2O = SO_{2}. It is also formed when substances containing sulphur are burned: ZnS + 3O = ZnO + SO_{2}. 2. _By the reduction of sulphuric acid. _ When concentrated sulphuricacid is heated with certain metals, such as copper, part of the acid ischanged into copper sulphate, and part is reduced to sulphurous acid. The latter then decomposes into sulphur dioxide and water, the completeequation being Cu + 2H_{2}SO_{4} = CuSO_{4} + SO_{2} + 2H_{2}O. 3. _By the action of an acid on a sulphite. _ Sulphites are salts ofsulphurous acid (H_{2}SO_{3}). When a sulphite is treated with an acid, sulphurous acid is set free, and being very unstable, decomposes intowater and sulphur dioxide. These reactions are expressed in theequations Na_{2}SO_{3} + 2HCl = 2NaCl + H_{2}SO_{3}, H_{2}SO_{3} = H_{2}O + SO_{2}. ~Explanation of the reaction. ~ In this case we have two reversiblereactions depending on each other. In the first reaction, (1) Na_{2}SO_{3} + 2HCl 2NaCl + H_{2}SO_{3}, we should expect an equilibrium to result, for none of the foursubstances in the equation are insoluble or volatile when water ispresent to hold them in solution. But the quantity of the H_{2}SO_{3} isconstantly diminishing, owing to the fact that it decomposes, asrepresented in the equation (2) H_{2}SO_{3} H_{2}O + SO_{2}, and the sulphur dioxide, being a gas, escapes. No equilibrium cantherefore result, since the quantity of the sulphurous acid isconstantly being diminished because of the escape of sulphur dioxide. ~Physical properties. ~ Sulphur dioxide is a colorless gas, which atordinary temperatures is 2. 2 times as heavy as air. It has a peculiar, irritating odor. The gas is very soluble in water, one volume of waterdissolving eighty of the gas under standard conditions. It is easilycondensed to a colorless liquid, and can be purchased in this conditionstored in strong bottles, such as the one represented in Fig. 42. [Illustration: Fig. 42] ~Chemical properties. ~ Sulphur dioxide has a marked tendency to combinewith other substances, and is therefore an active substance chemically. It combines with oxygen gas, but not very easily. It can, however, takeoxygen away from some other substances, and is therefore a good reducingagent. Its most marked chemical property is its ability to combine withwater to form sulphurous acid (H_{2}SO_{3}). ~Sulphurous acid~ (H_{2}SO_{3}). When sulphur dioxide dissolves in waterit combines chemically with it to form sulphurous acid, an unstablesubstance having the formula H_{3}SO_{3}. It is impossible to preparethis acid in pure form, as it breaks down very easily into water andsulphur dioxide. The reaction is therefore reversible, and is expressedby the equation H_{2}O + SO_{2} H_{2}SO_{3}. Solutions of the acid in water have a number of interesting properties. 1. _Acid properties. _ The solution has all the properties typical of anacid. When neutralized by bases, sulphurous acid yields a series ofsalts called _sulphites_. 2. _Reducing properties. _ Solutions of sulphurous acid act as goodreducing agents. This is due to the fact that sulphurous acid has thepower of taking up oxygen from the air, or from substances rich inoxygen, and is changed by this reaction into sulphuric acid: H_{2}SO_{3} + O = H_{2}SO_{4}, H_{2}SO_{3} + H_{2}O_{2} = H_{2}S0_{4} + H_{2}O. 3. _Bleaching properties. _ Sulphurous acid has strong bleachingproperties, acting upon many colored substances in such a way as todestroy their color. It is on this account used to bleach paper, strawgoods, and even such foods as canned corn. 4. _Antiseptic properties. _ Sulphurous acid has marked antisepticproperties, and on this account has the power of arrestingfermentation. It is therefore used as a preservative. ~Salts of sulphurous acid, --sulphites. ~ The sulphites, like sulphurousacid, have the power of taking up oxygen very readily, and are goodreducing agents. On account of this tendency, commercial sulphites areoften contaminated with sulphates. A great deal of sodium sulphite isused in the bleaching industry, and as a reagent for softening paperpulp. ~Sulphur trioxide~ (SO_{3}). When sulphur dioxide and oxygen are heatedtogether at a rather high temperature, a small amount of sulphurtrioxide (SO_{3}) is formed, but the reaction is slow and incomplete. If, however, the heating takes place in the presence of very fineplatinum dust, the reaction is rapid and nearly complete. [Illustration: Fig. 43] ~ Experimental preparation of sulphur trioxide. ~ The experiment can be performed by the use of the apparatus shown in Fig. 43, the fine platinum being secured by moistening asbestos fiber with a solution of platinum chloride and igniting it in a flame. The fiber, covered with fine platinum, is placed in a tube of hard glass, which is then heated with a burner to about 350°, while sulphur dioxide and air are passed into the tube. Union takes place at once, and the strongly fuming sulphur trioxide escapes from the jet at the end of the tube, and may be condensed by surrounding the receiving tube with a freezing mixture. ~Properties of sulphur trioxide. ~ Sulphur trioxide is a colorless liquid, which solidifies at about 15° and boils at 46°. A trace of moisturecauses it to solidify into a mass of silky white crystals, somewhatresembling asbestos fiber in appearance. In contact with the air itfumes strongly, and when thrown upon water it dissolves with a hissingsound and the liberation of a great deal of heat. The product of thisreaction is sulphuric acid, so that sulphur trioxide is the anhydride ofthat acid: SO_{3} + H_{2}O = H_{2}SO_{4}. ~Catalysis. ~ It has been found that many chemical reactions, such as theunion of sulphur dioxide with oxygen, are much influenced by thepresence of substances which do not themselves seem to take a part inthe reaction, and are left apparently unchanged after it has ceased. These reactions go on very slowly under ordinary circumstances, but aregreatly hastened by the presence of the foreign substance. Substanceswhich hasten very slow reactions in this way are said to act ascatalytic agents or _catalyzers_, and the action is called _catalysis_. Just how the action is brought about is not well understood. DEFINITION: _A catalyzer is a substance which changes the velocity of areaction, but does not change its products. _ ~Examples of Catalysis. ~ We have already had several instances of suchaction. Oxygen and hydrogen combine with each other at ordinarytemperatures in the presence of platinum powder, while if no catalyticagent is present they do not combine in appreciable quantities until arather high temperature is reached. Potassium chlorate, when heated withmanganese dioxide, gives up its oxygen at a much lower temperature thanwhen heated alone. Hydrogen dioxide decomposes very rapidly whenpowdered manganese dioxide is sifted into its concentrated solution. On the other hand, the catalytic agent sometimes retards chemicalaction. For example, a solution of hydrogen dioxide decomposes moreslowly when it contains a little phosphoric acid than when perfectlypure. For this reason commercial hydrogen dioxide always containsphosphoric acid. Many reactions are brought about by the catalytic action of traces ofwater. For example, phosphorus will not burn in oxygen in the absence ofall moisture. Hydrochloric acid will not unite with ammonia if thereagents are perfectly dry. It is probable that many of the chemicaltransformations in physiological processes, such as digestion, areassisted by certain substances acting as catalytic agents. The principleof catalysis is therefore very important. ~Sulphuric acid~ (_oil of vitriol_) (H_{2}SO_{4}). Sulphuric acid is oneof the most important of all manufactured chemicals. Not only is it oneof the most common reagents in the laboratory, but enormous quantitiesof it are used in many of the industries, especially in the refining ofpetroleum, the manufacture of nitroglycerin, sodium carbonate, andfertilizers. ~Manufacture of sulphuric acid. ~ 1. _Contact process_. The reactionstaking place in this process are represented by the following equations: SO_{2} + O = SO_{3}, SO_{3} + H_{2}O = H_{2}SO_{4}. To bring about the first of these reactions rapidly, a catalyzer isemployed, and the process is carried out in the following way: Largeiron tubes are packed with some porous material, such as calcium andmagnesium sulphates, which contains a suitable catalytic substancescattered through it. The catalyzers most used are platinum powder, vanadium oxide, and iron oxide. Purified sulphur dioxide and air arepassed through the tubes, which are kept at a temperature of about 350°. Sulphur trioxide is formed, and as it issues from the tube it isabsorbed in water or dilute sulphuric acid. The process is continueduntil all the water in the absorbing vessel has been changed intosulphuric acid, so that a very concentrated acid is made in this way. Anexcess of the trioxide may dissolve in the strong sulphuric acid, forming what is known as _fuming sulphuric acid_. 2. _Chamber process. _ The method of manufacture exclusively employeduntil recent years, and still in very extensive use, is much morecomplicated. The reactions are quite involved, but the conversion ofwater, sulphur dioxide, and oxygen into sulphuric acid is accomplishedby the catalytic action of oxides of nitrogen. The reactions are broughtabout in large lead-lined chambers, into which oxides of nitrogen, sulphur dioxide, steam, and air are introduced in suitable proportions. ~Reactions of the chamber process. ~ In a very general way, the various reactions which take place in the lead chambers may be expressed in two equations. In the first reaction sulphur dioxide, nitrogen peroxide, steam, and oxygen unite, as shown in the equation (1) 2SO_{2} + 2NO_{2} + H_{2}O + O = 2SO_{2} (OH) (NO_{2}). The product formed in this reaction is called nitrosulphuric acid or "chamber crystals. " It actually separates on the walls of the chambers when the process is not working properly. Under normal conditions, it is decomposed as fast as it is formed by the action of excess of steam, as shown in the equation (2) 2SO_{2} (OH) (NO_{2}) + H_{2}O + O = 2H_{2}SO_{4} + 2NO_{2}. The nitrogen dioxide formed in this reaction can now enter into combination with a new quantity of sulphur dioxide, steam, and oxygen, and the series of reactions go on indefinitely. Many other reactions occur, but these two illustrate the principle of the process. The relation between sulphuric acid and nitrosulphuric acid can be seenby comparing their structural formulas: O= -OH O= -OH S S O= -OH O= -NO_{2} The latter may be regarded as derived from the former by thesubstitution of the nitro group (NO_{2}) for the hydroxyl group (OH). [Illustration: Fig. 44] ~The sulphuric acid plant. ~ Fig. 44 illustrates the simpler parts of aplant used in the manufacture of sulphuric acid by the chamber process. Sulphur or some sulphide, as FeS_{2}, is burned in furnace A. Theresulting sulphur dioxide, together with air and some nitrogen peroxide, are conducted into the large chambers, the capacity of each chamberbeing about 75, 000 cu. Ft. Steam is also admitted into these chambers atdifferent points. These compounds react to form sulphuric acid, according to the equations given above. The nitrogen left after thewithdrawal of the oxygen from the admitted air escapes through theGay-Lussac tower X. In order to prevent the escape of the oxides ofnitrogen regenerated in the reaction, the tower is filled with lumps ofcoke, over which trickles concentrated sulphuric acid admitted from Y. The nitrogen peroxide dissolves in the acid and the resulting solutioncollects in H. This is pumped into E, where it is mixed with diluteacid and allowed to trickle down through the chamber D (Glover tower), which is filled with some acid-resisting rock. Here the nitrogenperoxide is expelled from the solution by the action of the hot gasesentering from A, and together with them enters the first chamberagain. The acid from which the nitrogen peroxide is expelled collects inF. Theoretically, a small amount of nitrogen peroxide would suffice toprepare an unlimited amount of sulphuric acid; practically, some of itescapes, and this is replaced by small amounts admitted at B. The sulphuric acid so formed, together with the excess of condensedsteam, collect upon the floor of the chambers in the form of a liquidcontaining from 62% to 70% of sulphuric acid. The product is called_chamber acid_ and is quite impure; but for many purposes, such as themanufacture of fertilizers, it needs no further treatment. It can beconcentrated by boiling it in vessels made of iron or platinum, whichresist the action of the acid, nearly all the water boiling off. Pureconcentrated acid can be made best by the contact process, while thechamber process is cheaper for the dilute impure acid. ~Physical properties. ~ Sulphuric acid is a colorless, oily liquid, nearlytwice as heavy as water. The ordinary concentrated acid contains about2% of water, has a density of 1. 84, and boils at 338°. It is sometimescalled _oil of vitriol_, since it was formerly made by distilling asubstance called _green vitriol_. ~Chemical properties. ~ Sulphuric acid possesses chemical properties whichmake it one of the most important of chemical substances. 1. _Action as an acid. _ In dilute solution sulphuric acid acts as anyother acid, forming salts with oxides and hydroxides. 2. _Action as an oxidizing agent. _ Sulphuric acid contains a largepercentage of oxygen and is, like nitric acid, a very good oxidizingagent. When the concentrated acid is heated with sulphur, carbon, andmany other substances, oxidation takes place, the sulphuric aciddecomposing according to the equation H_{2}SO_{4} = H_{2}SO_{3} + O. 3. _Action on metals. _ In dilute solution sulphuric acid acts upon manymetals, such as zinc, forming a sulphate and liberating hydrogen. Whenthe concentrated acid is employed the hydrogen set free is oxidized by anew portion of the acid, with the liberation of sulphur dioxide. Withcopper the reactions are expressed by the equations (1) Cu + H_{2}SO_{4} = CuSO_{4} + 2H, (2) H_{2}SO_{4} + 2H = H_{2}SO_{3} + H_{2}O, (3) H_{2}SO_{3} = H_{2}O + SO_{2}. By combining these equations the following one is obtained: Cu + 2H_{2}SO_{4} = CuSO_{4} + SO_{2} + 2H_{2}O. 4. _Action on salts. _ We have repeatedly seen that an acid of highboiling point heated with the salt of some acid of lower boiling pointwill drive out the low boiling acid. The boiling point of sulphuric acid(338°) is higher than that of almost any common acid; hence it is usedlargely in the preparation of other acids. 5. _Action on water. _ Concentrated sulphuric acid has a very greataffinity for water, and is therefore an effective dehydrating agent. Gases which have no chemical action upon sulphuric acid can be freedfrom water vapor by bubbling them through the strong acid. When the acidis diluted with water much heat is set free, and care must be taken tokeep the liquid thoroughly stirred during the mixing, and to pour theacid into the water, --never the reverse. Not only can sulphuric acid absorb water, but it will often withdraw theelements hydrogen and oxygen from a compound containing them, decomposing the compound, and combining with the water so formed. Forthis reason most organic substances, such as sugar, wood, cotton, andwoolen fiber, and even flesh, all of which contain much oxygen andhydrogen in addition to carbon, are charred or burned by the action ofthe concentrated acid. ~Salts of sulphuric acid, --sulphates. ~ The sulphates form a very importantclass of salts, and many of them have commercial uses. Copperas (ironsulphate), blue vitriol (copper sulphate), and Epsom salt (magnesiumsulphate) serve as examples. Many sulphates are important minerals, prominent among these being gypsum (calcium sulphate) and barytes(barium sulphate). ~Thiosulphuric acid~ (H_{2}S_{2}O_{3}); ~Thiosulphates. ~ Many other acids of sulphur containing oxygen are known, but none of them are of great importance. Most of them cannot be prepared in a pure state, and are known only through their salts. The most important of these is thiosulphuric acid. When sodium sulphite is boiled with sulphur the two substances combine, forming a salt which has the composition represented in the formula Na_{2}S_{2}O_{3}: Na_{2}SO_{3} + S = Na_{2}S_{2}O_{3}. The substance is called sodium thiosulphate, and is a salt of the easily decomposed acid H_{2}S_{2}O_{3}, called thiosulphuric acid. This reaction is quite similar to the action of oxygen upon sulphites: Na_{2}SO_{3} + O = Na_{2}SO_{4}. More commonly the salt is called sodium hyposulphite, or merely "hypo. " It is a white solid and is extensively used in photography, in the bleaching industry, and as a disinfectant. ~Monobasic and dibasic acids. ~ Such acids as hydrochloric and nitricacids, which have only one replaceable hydrogen atom in the molecule, orin other words yield one hydrogen ion in solution, are called monobasicacids. Acids yielding two hydrogen ions in solution are called dibasicacids. Similarly, we may have tribasic and tetrabasic acids. The threeacids of sulphur are dibasic acids. It is therefore possible for each ofthem to form both normal and acid salts. The acid salts can be made intwo ways: the acid may be treated with only half enough base toneutralize it, -- NaOH + H_{2}SO_{4} = NaHSO_{4} + H_{2}O; or a normal salt may be treated with the free acid, -- Na_{2}SO_{4} + H_{2}SO_{4} = 2NaHSO_{4}. Acid sulphites and sulphides may be made in the same ways. ~Carbon disulphide~ (CS_{2}). When sulphur vapor is passed over highlyheated carbon the two elements combine, forming carbon disulphide(CS_{2}), just as oxygen and carbon unite to form carbon dioxide(CO_{2}). The substance is a heavy, colorless liquid, possessing, whenpure, a pleasant ethereal odor. On standing for some time, especiallywhen exposed to sunlight, it undergoes a slight decomposition andacquires a most disagreeable, rancid odor. It has the property ofdissolving many substances, such as gums, resins, and waxes, which areinsoluble in most liquids, and it is extensively used as a solvent forsuch substances. It is also used as an insecticide. It boils at a lowtemperature (46°), and its vapor is very inflammable, burning in the airto form carbon dioxide and sulphur dioxide, according to the equation CS_{2} + 6O = CO_{2} + 2SO_{2}. [Illustration: Fig. 45] ~Commercial preparation of carbon disulphide. ~ In the preparation of carbon disulphide an electrical furnace is employed, such as is represented in Fig. 45. The furnace is packed with carbon C, and this is fed in through the hoppers B, as fast as that which is present in the hearth of the furnace is used up. Sulphur is introduced at A, and at the lower ends of the tubes it is melted by the heat of the furnace and flows into the hearth as a liquid. An electrical current is passed through the carbon and melted sulphur from the electrodes E, heating the charge. The vapors of carbon disulphide pass up through the furnace and escape at D, from which they pass to a suitable condensing apparatus. ~Comparison of sulphur and oxygen. ~ A comparison of the formulas and thechemical properties of corresponding compounds of oxygen and sulphurbrings to light many striking similarities. The conduct ofhydrosulphuric acid and water toward many substances has been seen to bevery similar; the oxides and sulphides of the metals have analogousformulas and undergo many parallel reactions. Carbon dioxide anddisulphide are prepared in similar ways and undergo many analogousreactions. It is clear, therefore, that these two elements are far moreclosely related to each other than to any of the other elements so farstudied. ~Selenium and tellurium. ~ These two very uncommon elements are still moreclosely related to sulphur than is oxygen. They occur in comparativelysmall quantities and are usually found associated with sulphur andsulphides, either as the free elements or more commonly in combinationwith metals. They form compounds with hydrogen of the formulas H_{2}Seand H_{2}Te; these bodies are gases with properties very similar tothose of H_{2}S. They also form oxides and oxygen acids which resemblethe corresponding sulphur compounds. The elements even have allotropicforms corresponding very closely to those of sulphur. Tellurium issometimes found in combination with gold and copper, and occasions somedifficulties in the refining of these metals. The elements have very fewpractical applications. ~Crystallography. ~ In order to understand the difference between the twokinds of sulphur crystals, it is necessary to know something aboutcrystals in general and the forms which they may assume. An examinationof a large number of crystals has shown that although they may differmuch in geometric form, they can all be considered as modifications of afew simple plans. The best way to understand the relation of one crystalto another is to look upon every crystal as having its faces and anglesarranged in definite fashion about certain imaginary lines drawnthrough the crystal. These lines are called axes, and bear much the samerelation to a crystal as do the axis and parallels of latitude andlongitude to the earth and a geographical study of it. All crystals canbe referred to one of six simple plans or systems, which have their axesas shown in the following drawings. The names and characteristics of these systems are as follows: 1. Isometric or regular system (Fig. 46). Three equal axes, all at rightangles. [Illustration: Fig. 46] 2. Tetragonal system (Fig. 47). Two equal axes and one of differentlength, all at right angles to each other. [Illustration: Fig. 47] 3. Orthorhombic system (Fig. 48). Three unequal axes, all at rightangles to each other. [Illustration: Fig. 48] 4. Monoclinic system (Fig. 49). Two axes at right angles, and a third atright angles to one of these, but inclined to the other. [Illustration: Fig. 49] 5. Triclinic system (Fig. 50). Three axes, all inclined to each other. [Illustration: Fig. 50] 6. Hexagonal system (Fig. 51). Three equal axes in the same planeintersecting at angles of 60°, and a fourth at right angles to all ofthese. [Illustration: Fig. 51] Every crystal can be imagined to have its faces and angles arranged in adefinite way around one of these systems of axes. A cube, for instance, is referred to Plan 1, an axis ending in the center of each face; whilein a regular octohedron an axis ends in each solid angle. These formsare shown in Fig. 46. It will be seen that both of these figures belongto the same system, though they are very different in appearance. In thesame way, many geometric forms may be derived from each of the systems, and the light lines about the axes in the drawings show two of thesimplest forms of each of the systems. In general a given substance always crystallizes in the same system, andtwo corresponding faces of each crystal of it always make the same anglewith each other. A few substances, of which sulphur is an example, crystallize in two different systems, and the crystals differ in suchphysical properties as melting point and density. Such substances aresaid to be _dimorphous_. EXERCISES 1. (a) Would the same amount of heat be generated by the combustion of1 g. Of each of the allotropic modifications of sulphur? (b) Would thesame amount of sulphur dioxide be formed in each case? 2. Is the equation for the preparation of hydrosulphuric acid areversible one? As ordinarily carried out, does the reaction completeitself? 3. Suppose that hydrosulphuric acid were a liquid, would it be necessaryto modify the method of preparation? 4. Can sulphuric acid be used to dry hydrosulphuric acid? Give reasonfor answer. 5. Does dry hydrosulphuric acid react with litmus paper? State reasonfor answer. 6. How many grams of iron sulphide are necessary to prepare 100 l. Ofhydrosulphuric acid when the laboratory conditions are 17° and 740 mm. Pressure? 7. Suppose that the hydrogen in 1 l. Of hydrosulphuric acid wereliberated; what volume would it occupy, the gases being measured underthe same conditions? 8. Write the equations representing the reaction between hydrosulphuricacid and sodium hydroxide and ammonium hydroxide respectively. 9. Show that the preparation of sulphur dioxide from a sulphite issimilar in principle to the preparation of hydrogen sulphide. 10. (a) Does dry sulphur dioxide react with litmus paper? (b) Howcan it be shown that a solution of sulphur dioxide in water acts like anacid? 11. (a) Calculate the percentage composition of sulphurous anhydrideand sulphuric anhydride. (b) Show how these two substances are inharmony with the law of multiple proportion. 12. How many pounds of sulphur would be necessary in the preparation of100 lb. Of 98% sulphuric acid? 13. What weight of sulphur dioxide is necessary in the preparation of 1kg. Of sodium sulphite? 14. What weight of copper sulphate crystals can be obtained bydissolving 1 kg. Of copper in sulphuric acid and crystallizing theproduct from water? 15. Write the names and formulas of the oxides and oxygen acids ofselenium and tellurium. 16. In the commercial preparation of carbon disulphide, what is thefunction of the electric current? 17. If the Gay-Lussac tower were omitted from the sulphuric acidfactory, what effect would this have on the cost of production ofsulphuric acid? CHAPTER XV PERIODIC LAW A number of the elements have now been studied somewhat closely. Thefirst three of these, oxygen, hydrogen, and nitrogen, while having somephysical properties in common with each other, have almost no point ofsimilarity as regards their chemical conduct. On the other hand, oxygenand sulphur, while quite different physically, have much in common intheir chemical properties. About eighty elements are now known. If all of these should haveproperties as diverse as do oxygen, hydrogen, and nitrogen, the study ofchemistry would plainly be a very difficult and complicated one. If, however, the elements can be classified in groups, the members of whichhave very similar properties, the study will be very much simplified. ~Earlier classification of the elements. ~ Even at an early period effortswere made to discover some natural principle in accordance with whichthe elements could be classified. Two of these classifications may bementioned here. 1. _Classification into metals and non-metals. _ The classification intometals and non-metals most naturally suggested itself. This grouping wasbased largely on physical properties, the metals being heavy, lustrous, malleable, ductile, and good conductors of heat and electricity. Elements possessing these properties are usually base-forming incharacter, and the ability to form bases came to be regarded as acharacteristic property of the metals. The non-metals possessedphysical properties which were the reverse of those of the metals, andwere acid-forming in character. Not much was gained by this classification, and it was very imperfect. Some metals, such as potassium, are very light; some non-metals, such asiodine, have a high luster; some elements can form either an acid or abase. 2. _Classification into triad families. _ In 1825 Döbereiner observedthat an interesting relation exists between the atomic weights ofchemically similar elements. To illustrate, lithium, sodium, andpotassium resemble each other very closely, and the atomic weight ofsodium is almost exactly an arithmetical mean between those of the othertwo: (7. 03 + 39. 15)/2 = 23. 09. In many chemical and physical propertiessodium is midway between the other two. A number of triad families were found, but among eighty elements, whoseatomic weights range all the way from 1 to 240, such agreements might bemere chance. Moreover many elements did not appear to belong to suchfamilies. ~Periodic division. ~ In 1869 the Russian chemist Mendeléeff devised anarrangement of the elements based on their atomic weights, which hasproved to be of great service in the comparative study of the elements. A few months later the German, Lothar Meyer, independently suggested thesame ideas. This arrangement brought to light a great generalization, now known as the _periodic law_. An exact statement of the law will begiven after the method of arranging the elements has been described. [Illustration: DMITRI IVANOVITCH MENDELÉEFF (Russian) (1834-1907) Author of the periodic law; made many investigations on the physicalconstants of elements and compounds; wrote an important book entitled"Principles of Chemistry"; university professor and governmentofficial] ~Arrangement of the periodic table. ~ The arrangement suggested byMendeléeff, modified somewhat by more recent investigations, is asfollows: Beginning with lithium, which has an atomic weight of 7, theelements are arranged in a horizontal row in the order of their atomicweights, thus: ~Li (7. 03), Be (9. 1), B (11), C (12), N (14. 04), O (16), F (19). ~ These seven elements all differ markedly from each other. The eighthelement, sodium, is very similar to lithium. It is placed just underlithium, and a new row follows: ~Na(23. 05), Mg (24. 36), Al (27. 1), Si (28. 4), P (31), S (32. 06), Cl(35. 45). ~ When the fifteenth element, potassium, is reached, it is placed undersodium, to which it is very similar, and serves to begin a third row: ~K (39. 15), Ca (40. 1), Sc (44. 1, ) Ti (48. 1), V (51. 2), Cr (52. 1), Mn(55). ~ Not only is there a strong similarity between lithium, sodium, andpotassium, which have been placed in a vertical row because of thisresemblance, but the elements in the other vertical rows exhibit much ofthe same kind of similarity among themselves, and evidently form littlenatural groups. The three elements following manganese, namely, iron, nickel, andcobalt, have atomic weights near together, and are very similarchemically. They do not strongly resemble any of the elements so farconsidered, and are accordingly placed in a group by themselves, following manganese. A new row is begun with copper, which somewhatresembles the elements of the first vertical column. Following the fifthand seventh rows are groups of three closely related elements, so thatthe completed arrangement has the appearance represented in the table onpage 168. THE PERIODIC ARRANGEMENT OF THE ELEMENTS --------+-----------+-----------+-----------+-----------+-----------+Periods | GROUP | GROUP | GROUP | GROUP | GROUP | | 0 | I | II | III | IV | |A B|A B|A B|A B|A B|--------+-----------+-----------+-----------+-----------+-----------+1 |H==1. 008 | | | | |2 |He=4 |Li=7. 03 |Be=9. 1 |B=11 |C=12 |--------+-----------+-----------+-----------+-----------+-----------+3 | Ne=20|Na=23. 05 | Mg=24. 36| AL=27. 1| Si=28. 4|--------+-----------+-----------+-----------+-----------+-----------+4 |A=39. 9 |K=39. 15 |Ca=40. 1 |Sc=44. 1 |Ti=48. 1 | | | | | | | | | | | | |--------+-----------+-----------+-----------+-----------+-----------+5 | | Cu=63. 6| Zn=65. 4| Ga=70| Ge=72. 5|--------+-----------+-----------+-----------+-----------+-----------+6 |Kr=81. 8 |Rb=85. 5 |Sr=87. 6 |Y=89 |Zr=90. 6 | | | | | | | | | | | | |--------+-----------+-----------+-----------+-----------+-----------+7 | | Ag=107. 93| Cd=112. 4| In=115| Sn=119|--------+-----------+-----------+-----------+-----------+-----------+8 |X=128 |Cs=132. 9 |Ba=137. 4 |La=138. 9 |Ce=Yb* | | | | | |140. 25-173 | | | | | | |--------+-----------+-----------+-----------+-----------+-----------+9 | Au=197. 2| Hg=200| Tl=204. 1| Pb=206. 9| Bi=208. 5|--------+-----------+-----------+-----------+-----------+-----------+10 | | |Ra=225 | |Th=232. 5 |--------+-----------+-----------+-----------+-----------+-----------+ | | R_{2}O | RO |R_{2}O_{3} | RO_{2} | | | RH | RH_{2} | RH_{3} | RH_{4} |--------+-----------+-----------+-----------+-----------+-----------+ ==================part 2============== --------+-----------+-----------+-----------+-----------+Periods | GROUP | GROUP | GROUP | GROUP | | V | VI | VII | VIII | |A B|A B|A B| |--------+-----------+-----------+-----------+-----------+1 | | | | |2 |N=14. 04 |O=16 |F=19 | |--------+-----------+-----------+-----------+-----------+3 | P=31| S=32. 06| Cl=35. 45| |--------+-----------+-----------+-----------+-----------+4 |V=51. 2 |Cr=52. 1 |Mn=55 |Fe=55. 9 | | | | |Ni=58. 7 | | | | |Co=59 |--------+-----------+-----------+-----------+-----------+5 | As=75| Se=79. 2| Br=79. 96| |--------+-----------+-----------+-----------+-----------+6 |Cb=94 |Mo=96 | |Ru=101. 7 | | | | |Rh=103 | | | | |Pd=106. 5 |--------+-----------+-----------+-----------+-----------+7 | Sb=120. 2| Te=127. 6| I=126. 97| |--------+-----------+-----------+-----------+-----------+8 |Ta=183 |W=184 | |Os=191 | | | | |Ir=193 | | | | |Pt=194. 8 |--------+-----------+-----------+-----------+-----------+9 | | | | |--------+-----------+-----------+-----------+-----------+10 | U=238. 5 | | | |--------+-----------+-----------+-----------+-----------+ | R_{2}O_{5}| RO_{3} | R_{2}O_{7}| RO_{4} | | RH_{3} | RH_{2} | RH | |--------+-----------+-----------+-----------+-----------+ [* This includes a number of elements whose atomic weights liebetween 140 and 173, but which have not been accurately studied, andso their proper arrangement is uncertain. ] ~Place of the atmospheric elements. ~ When argon was discovered it was seenat once that there was no place in the table for an element of atomicweight approximately 40. When the other inactive elements were found, however, it became apparent that they form a group just preceding Group1. They are accordingly arranged in this way in Group 0 (see table onopposite page). A study of this table brings to light certain verystriking facts. ~Properties of elements vary with atomic weights. ~ There is evidently aclose relation between the properties of an element and its atomicweight. Lithium, at the beginning of the first group, is a very strongbase-forming element, with pronounced metallic properties. Beryllium, following lithium, is less strongly base-forming, while boron has somebase-forming and some acid-forming properties. In carbon allbase-forming properties have disappeared, and the acid-formingproperties are more marked than in boron. These become still moreemphasized as we pass through nitrogen and oxygen, until on reachingfluorine we have one of the strongest acid-forming elements. Theproperties of these seven elements therefore vary regularly with theiratomic weights, or, in mathematical language, are regular functions ofthem. ~Periodic law. ~ The properties of the first seven elements vary_continuously_--that is steadily--away from base-forming and towardacid-forming properties. If lithium had the smallest atomic weight ofany of the elements, and fluorine the greatest, so that in passing fromone to the other we had included all the elements, we could say that theproperties of elements are continuous functions of their atomic weights. But fluorine is an element of small atomic weight, and the one followingit, sodium, breaks the regular order, for in it reappear all thecharacteristic properties of lithium. Magnesium, following sodium, bearsmuch the same relation to beryllium that sodium does to lithium, andthe properties of the elements in the second row vary much as they do inthe first row until potassium is reached, when another repetitionbegins. The properties of the elements do not vary continuously, therefore, with atomic weights, but at regular intervals there is arepetition, or _period_. This generalization is known as the _periodiclaw_, and may be stated thus: _The properties of elements are periodicfunctions of their atomic weights. _ ~The two families in a group. ~ While all the elements in a given verticalcolumn bear a general resemblance to each other, it has been noticedthat those belonging to periods having even numbers are very strikinglysimilar to each other. They are placed at the left side of the groupcolumns. In like manner, the elements belonging to the odd periods arevery similar and are arranged at the right side of the group columns. Thus calcium, strontium, and barium are very much alike; so, too, aremagnesium, zinc, and cadmium. The resemblance between calcium andmagnesium, or strontium and zinc, is much less marked. This method ofarrangement therefore divides each group into two families, eachcontaining four or five members, between which there is a greatsimilarity. ~Family resemblances. ~ Let us now inquire more closely in what respectsthe elements of a family resemble each other. 1. _Valence. _ In general the valence of the elements in a family is thesame, and the formulas of their compounds are therefore similar. If weknow that the formula of sodium chloride is NaCl, it is pretty certainthat the formula of potassium chloride will be KCl--not KCl_{2} orKCl_{3}. The general formulas R_{2}O, RO, etc. , placed below thecolumns show the formulas of the oxides of the elements in the columnprovided they form oxides. In like manner the formulas RH, RH_{2}, etc. , show the composition of the compounds formed with hydrogen or chlorine. 2. _Chemical properties. _ The chemical properties of the members of afamily are quite similar. If one member is a metal, the others usuallyare; if one is a non-metal, so, too, are the others. The families in thefirst two columns consist of metals, while the elements found in thelast two columns form acids. There is in addition a certain regularityin properties of the elements in each family. If the element at the headof the family is a strong acid-forming element, this property is likelyto diminish gradually, as we pass to the members of the family withhigher atomic weights. Thus phosphorus is strongly acid-forming, arsenicless so, antimony still less so, while bismuth has almost noacid-forming properties. We shall meet with many illustrations of thisfact. 3. _Physical properties. _ In the same way, the physical properties ofthe members of a family are in general somewhat similar, and show aregular gradation as we pass from element to element in the family. Thusthe densities of the members of the magnesium family are Mg = 1. 75, Zn = 7. 00, Cd = 8. 67, Hg = 13. 6. Their melting points are Mg = 750°, Zn = 420°, Cd = 320°, Hg = -39. 5°. ~Value of the periodic law. ~ The periodic law has proved of much value inthe development of the science of chemistry. 1. _It simplifies study. _ It is at once evident that such regularitiesvery much simplify the study of chemistry. A thorough study of oneelement of a family makes the study of the other members a much easiertask, since so many of the properties and chemical reactions of theelements are similar. Thus, having studied the element sulphur in somedetail, it is not necessary to study selenium and tellurium so closely, for most of their properties can be predicted from the relation whichthey sustain to sulphur. 2. _It predicts new elements. _ When the periodic law was firstformulated there were a number of vacant places in the table whichevidently belonged to elements at that time unknown. From their positionin the table, Mendeléeff predicted with great precision the propertiesof the elements which he felt sure would one day be discovered to fillthese places. Three of them, scandium, germanium, and gallium, werefound within fifteen years, and their properties agreed in a remarkableway with the predictions of Mendeléeff. There are still some vacantplaces in the table, especially among the heavier elements. 3. _It corrects errors. _ The physical constants of many of the elementsdid not at first agree with those demanded by the periodic law, and afurther study of many such cases showed that errors had been made. Thelaw has therefore done much service in indicating probable error. ~Imperfections of the law. ~ There still remain a good many features whichmust be regarded as imperfections in the law. Most conspicuous is thefact that the element hydrogen has no place in the table. In some of thegroups elements appear in one of the families, while all of theirproperties show that they belong in the other. Thus sodium belongs withlithium and not with copper; fluorine belongs with chlorine and not withmanganese. There are two instances where the elements must betransposed in order to make them fit into their proper group. Accordingto their atomic weights, tellurium should follow iodine, and argonshould follow potassium. Their properties show in each case that thisorder must be reversed. The table separates some elements altogetherwhich, in many respects have closely agreeing properties. Iron, chromium, and manganese are all in different groups, although they aresimilar in many respects. The system is therefore to be regarded as but a partial and imperfectexpression of some very important and fundamental relation between thesubstances which we know as elements, the exact nature of this relationbeing as yet not completely clear to us. EXERCISES 1. Suppose that an element were discovered that filled the blank inGroup O, Period 5; what properties would it probably have? 2. Suppose that an element were discovered that filled the blank inGroup VI, Period 9, family B; what properties would it have? 3. Sulphur and oxygen both belong in Group VI, although in differentfamilies; in what respects are the two similar? CHAPTER XVI THE CHLORINE FAMILY ================================================================== | | | | | ATOMIC | MELTING | BOILING | COLOR AND STATE | WEIGHT | POINT | POINT |______________|________|_________|_________|______________________ | | | |Fluorine (F) | 19. 00 | -223° | -187° | Pale yellowish gas. Chlorine (Cl) | 35. 45 | -102° | -33. 6° | Greenish-yellow gas. Bromine (Br) | 79. 96 | -7° | 59° | Red liquid. Iodine (I) | 126. 97 | 107° | 175° | Purplish-black solid. ================================================================== ~The family. ~ The four elements named in the above table form a stronglymarked family of elements and illustrate very clearly the way in whichthe members of a family in a periodic group resemble each other, as wellas the character of the differences which we may expect to find betweenthe individual members. 1. _Occurrence. _ These elements do not occur in nature in the freestate. The compounds of the last three elements of the family are foundextensively in sea water, and on this account the name _halogens_, signifying "producers of sea salt, " is sometimes applied to the family. 2. _Properties. _ As will be seen by reference to the table, the meltingpoints and boiling points of the elements of the family increase withtheir atomic weights. A somewhat similar gradation is noted in theircolor and state. One atom of each of the elements combines with one atomof hydrogen to form acids, which are gases very soluble in water. Theaffinity of the elements for hydrogen is in the inverse order of theiratomic weights, fluorine having the strongest affinity and iodine theweakest. Only chlorine and iodine form oxides, and those of the formerelement are very unstable. The elements of the group are univalent intheir compounds with hydrogen and the metals. FLUORINE ~Occurrence. ~ The element fluorine occurs in nature most abundantly as themineral fluorspar (CaF_{2}), as cryolite (Na_{3}AlF_{6}), and in thecomplex mineral apatite (3 Ca_{3}(PO_{4})_{2}·CaF_{2}). ~Preparation. ~ All attempts to isolate the element resulted in failureuntil recent years. Methods similar to those which succeed in thepreparation of the other elements of the family cannot be used; for assoon as the fluorine is liberated it combines with the materials ofwhich the apparatus is made or with the hydrogen of the water which isalways present. The preparation of fluorine was finally accomplished bythe French chemist Moissan by the electrolysis of hydrofluoric acid. Perfectly dry hydrofluoric acid (HF) was condensed to a liquid andplaced in a U-shaped tube made of platinum (or copper), which wasfurnished with electrodes and delivery tubes, as shown in Fig. 52. Thisliquid is not an electrolyte, but becomes such when potassium fluorideis dissolved in it. When this solution was electrolyzed hydrogen was setfree at the cathode and fluorine at the anode. [Illustration: Fig. 52] ~Properties. ~ Fluorine is a gas of slightly yellowish color, and can becondensed to a liquid boiling at -187° under atmospheric pressure. Itsolidifies at -223°. It is extremely active chemically, being the mostactive of all the elements at ordinary temperatures. It combines with all the common elements save oxygen, very often withincandescence and the liberation of much heat. It has a strong affinityfor hydrogen and is able to withdraw it from its compounds with otherelements. Because of its great activity it is extremely poisonous. Fluorine does not form any oxides, neither does it form any oxygenacids, in which respects it differs from the other members of thefamily. ~Hydrofluoric acid~ (HF). Hydrofluoric acid is readily obtained fromfluorspar by the action of concentrated sulphuric acid. The equation is CaF_{2} + H_{2}SO_{4} = CaSO_{4} + 2HF. In its physical properties it resembles the binary acids of the otherelements of this family, being, however, more easily condensed to aliquid. The anhydrous acid boils at 19° and can therefore be prepared atordinary pressures. It is soluble in all proportions in water, and aconcentrated solution--about 50%--is prepared for the market. Its fumesare exceedingly irritating to the respiratory organs, and severalchemists have lost their lives by accidentally breathing them. [Illustration: HENRI MOISSAN (French) (1853-1907) Famous for his work with the electric furnace at high temperatures;prepared artificial diamonds, together with many new binary compoundssuch as carbides, silicides, borides, and nitrides; isolated fluorineand studied its properties and its compounds very thoroughly] ~Chemical properties. ~ Hydrofluoric acid, like other strong acids, readilyacts on bases and metallic oxides and forms the corresponding fluorides. It also dissolves certain metals such as silver and copper. It acts veryvigorously upon organic matter, a single drop of the concentrated acidmaking a sore on the skin which is very painful and slow in healing. Itsmost characteristic property is its action upon silicon dioxide(SiO_{2}), with which it forms water and the gas silicon tetrafluoride(SiF_{4}), as shown in the equation SiO_{2} + 4HF = SiF_{4} + 2H_{2}O. Glass consists of certain compounds of silicon, which are likewise actedon by the acid so that it cannot be kept in glass bottles. It ispreserved in flasks made of wax or gutta-percha. ~Etching. ~ Advantage is taken of this reaction in etching designs upon glass. The glass vessel is painted over with a protective paint upon which the acid will not act, the parts which it is desired to make opaque being left unprotected. A mixture of fluorspar and sulphuric acid is then painted over the vessel and after a few minutes the vessel is washed clean. Wherever the hydrofluoric acid comes in contact with the glass it acts upon it, destroying its luster and making it opaque, so that the exposed design will be etched upon the clear glass. Frosted glass globes are often made in this way. The etching may also be effected by covering the glass with a thin layer of paraffin, cutting the design through the wax and then exposing the glass to the fumes of the acid. ~Salts of hydrofluoric acid, --fluorides. ~ A number of the fluorides areknown, but only one of them, calcium fluoride (CaF_{2}), is ofimportance. This is the well-known mineral fluorspar. CHLORINE ~Historical. ~ While studying the action of hydrochloric acid upon themineral pyrolusite, in 1774, Scheele obtained a yellowish, gaseoussubstance to which he gave a name in keeping with the phlogiston theorythen current. Later it was supposed to be a compound containing oxygen. In 1810, however, the English chemist Sir Humphry Davy proved it to bean element and named it chlorine. ~Occurrence. ~ Chlorine does not occur free in nature, but its compoundsare widely distributed. For the most part it occurs in combination withthe metals in the form of chlorides, those of sodium, potassium, andmagnesium being most abundant. Nearly all salt water contains thesesubstances, particularly sodium chloride, and very large salt bedsconsisting of chlorides are found in many parts of the world. ~Preparation. ~ Two general methods of preparing chlorine may be mentioned, namely, the laboratory method and the electrolytic method. 1. _Laboratory method. _ In the laboratory chlorine is made by warmingthe mineral pyrolusite (manganese dioxide, MnO_{2}) with concentratedhydrochloric acid. The first reaction, which seems to be similar to theaction of acids upon oxides in general, is expressed in the equation MnO_{2} + 4HCl = MnCl_{4} + 2H_{2}O. The manganese compound so formed is very unstable, however, and breaksclown according to the equation MnCl_{4} = MnCl_{2} + 2Cl. Instead of using hydrochloric acid in the preparation of chlorine itwill serve just as well to use a mixture of sodium chloride andsulphuric acid, since these two react to form hydrochloric acid. Thefollowing equations will then express the changes: (1) 2NaCl + H_{2}SO_{4} = Na_{2}SO_{4} + 2HCl. (2) MnO_{2} + 4 HCl = MnCl_{2} + 2Cl + 2H_{2}O. (3) MnCl_{2} + H_{2}SO_{4} = MnSO_{4} + 2HCl. Combining these equations, the following equation expressing thecomplete reaction is obtained: 2NaCl + MnO_{2} + 2H_{2}SO_{4} = MnSO_{4} + Na_{2}SO_{4} + 2H_{2}O + 2Cl. Since the hydrochloric acid liberated in the third equation is free toact upon manganese dioxide, it will be seen that all of the chlorineoriginally present in the sodium chloride is set free. The manganese dioxide and the hydrochloric acid are brought together in a flask, as represented in Fig. 53, and a gentle heat is applied. The rate of evolution of the gas is regulated by the amount of heat applied, and the gas is collected by displacement of air. As the equations show, only half of the chlorine present in the hydrochloric acid is liberated. [Illustration: Fig. 53] 2. _Electrolytic method. _ Under the discussion of electrolysis (p. 102)it was shown that when a solution of sodium chloride is electrolyzedchlorine is evolved at the anode, while the sodium set free at thecathode reacts with the water to form hydrogen, which is evolved, andsodium hydroxide, which remains in solution. A great deal of thechlorine required in the chemical industries is now made in this way inconnection with the manufacture of sodium hydroxide. ~Physical properties. ~ Chlorine is a greenish-yellow gas, which has apeculiar suffocating odor and produces a very violent effect upon thethroat and lungs. Even when inhaled in small quantities it oftenproduces all the symptoms of a hard cold, and in larger quantities mayhave serious and even fatal action. It is quite heavy (density = 2. 45)and can therefore be collected by displacement of air. One volume ofwater under ordinary conditions dissolves about three volumes ofchlorine. The gas is readily liquefied, a pressure of six atmospheresserving to liquefy it at 0°. It forms a yellowish liquid whichsolidifies at -102°. ~Chemical properties. ~ At ordinary temperatures chlorine is far moreactive chemically than any of the elements we have so far considered, with the exception of fluorine; indeed, it is one of the most active ofall elements. 1. _Action on metals. _ A great many metals combine directly withchlorine, especially when hot. A strip of copper foil heated in a burnerflame and then dropped into chlorine burns with incandescence. Sodiumburns brilliantly when heated strongly in slightly moist chlorine. Goldand silver are quickly tarnished by the gas. 2. _Action on non-metals. _ Chlorine has likewise a strong affinity formany of the non-metals. Thus phosphorus burns in a current of the gas, while antimony and arsenic in the form of a fine powder at once burstinto flame when dropped into jars of the gas. The products formed in allcases where chlorine combines with another element are called_chlorides_. 3. _Action on hydrogen. _ Chlorine has a strong affinity for hydrogen, uniting with it to form hydrochloric acid. A jet of hydrogen burning inthe air continues to burn when introduced into a jar of chlorine, givinga somewhat luminous flame. A mixture of the two gases explodes violentlywhen a spark is passed through it or when it is exposed to brightsunlight. In the latter case it is the light and not the heat whichstarts the action. 4. _Action on substances containing hydrogen. _ Not only will chlorinecombine directly with free hydrogen but it will often abstract theelement from its compounds. Thus, when chlorine is passed into asolution containing hydrosulphuric acid, sulphur is precipitated andHydrochloric acid formed. The reaction is shown by the followingequation: H_{2}S + 2Cl = 2HCl + S. With ammonia the action is similar: NH_{3} + 3Cl = 3HCl + N. The same tendency is very strikingly seen in the action of chlorine uponturpentine. The latter substance is largely made up of compounds havingthe composition represented by the formula C_{10}H_{16}. When a strip ofpaper moistened with warm turpentine is placed in a jar of chlorinedense fumes of hydrochloric acid appear and a black deposit of carbon isformed. Even water, which is a very stable compound, can be decomposedby chlorine, the oxygen being liberated. This may be shown in thefollowing way: [Illustration: Fig. 54] If a long tube of rather large diameter is filled with a strong solution of chlorine in water and inverted in a vessel of the same solution, as shown in Fig. 54, and the apparatus is placed in bright sunlight, very soon bubbles of a gas will be observed to rise through the solution and collect in the tube. An examination of this gas will show that it is oxygen. It is liberated from water in accordance with the following equation: H_{2}O + 2Cl = 2HCl + O. 5. _Action on color substances, --bleaching action. _ If strips ofbrightly colored cloth or some highly colored flowers are placed inquite dry chlorine, no marked change in color is noticed as a rule. If, however, the cloth and flowers are first moistened, the color rapidlydisappears, that is, the objects are bleached. Evidently the moisture aswell as the chlorine is concerned in the action, and a study of the caseshows that the chlorine has combined with the hydrogen of the water. Theoxygen set free oxidizes the color substance, converting it into acolorless compound. It is evident from this explanation that chlorinewill only bleach those substances which are changed into colorlesscompounds by oxidation. 6. _Action as a disinfectant. _ Chlorine has also marked germicidalproperties, and the free element, as well as compounds from which it iseasily liberated, are used as disinfectants. ~Nascent state. ~ It will be noticed that oxygen when set free from waterby chlorine is able to do what ordinary oxygen cannot do, for both thecloth and the flowers are unchanged in the air which contains oxygen. Itis generally true that the activity of an element is greatest at theinstant of liberation from its compounds. To express this fact elementsat the instant of liberation are said to be in the _nascent state_. Itis nascent oxygen which does the bleaching. ~Hydrochloric acid~ (_muriatic acid_) (HCl). The preparation ofhydrochloric acid may be discussed under two general heads: 1. _Laboratory preparation. _ The product formed by the burning ofhydrogen in chlorine is the gas hydrochloric acid. This substance ismuch more easily obtained, however, by treating common salt (sodiumchloride) with sulphuric acid. The following equation shows thereaction: 2NaCl + H_{2}SO_{4} = Na_{2}SO_{4} + 2HCl. The dry salt is placed in a flask furnished with a funnel tube and anexit tube, the sulphuric acid is added, and the flask gently warmed. Thehydrochloric acid gas is rapidly given off and can be collected bydisplacement of air. The same apparatus can be used as was employed inthe preparation of chlorine (Fig. 53). When a _solution_ of salt is treated with sulphuric acid there is no very marked action. The hydrochloric acid formed is very soluble in water, and so does not escape from the solution; hence a state of equilibrium is soon reached between the four substances represented in the equation. When _concentrated_ sulphuric acid, in which hydrochloric acid is not soluble, is poured upon dry salt the reaction is complete. 2. _Commercial preparation. _ Commercially, hydrochloric acid is preparedin connection with the manufacture of sodium sulphate, the reactionbeing the same as that just given. The reaction is carried out in afurnace, and the hydrochloric acid as it escapes in the form of gas ispassed into water in which it dissolves, the solution forming thehydrochloric acid of commerce. When the materials are pure a colorlesssolution is obtained. The most concentrated solution has a density of1. 2 and contains 40% HCl. The commercial acid, often called _muriaticacid_, is usually colored yellow by impurities. ~Composition of hydrochloric acid. ~ When a solution of hydrochloric acidis electrolyzed in an apparatus similar to the one in which water waselectrolyzed (Fig. 18), chlorine collects at the anode and hydrogen atthe cathode. At first the chlorine dissolves in the water, but soon thewater in the one tube becomes saturated with it, and if the stopcocksare left open until this is the case, and are then closed, it will beseen that the two gases are set free in equal volumes. When measured volumes of the two gases are caused to unite it is foundthat one volume of hydrogen combines with one of chlorine. Otherexperiments show that the volume of hydrochloric acid formed is justequal to the sum of the volumes of hydrogen and chlorine. Therefore onevolume of hydrogen combines with one volume of chlorine to form twovolumes of hydrochloric acid gas. Since chlorine is 35. 18 times as heavyas hydrogen, it follows that one part of hydrogen by weight combineswith 35. 18 parts of chlorine to form 36. 18 parts of hydrochloric acid. ~Physical properties. ~ Hydrochloric acid is a colorless gas which has anirritating effect when inhaled, and possesses a sour, biting taste, butno marked odor. It is heavier than air (density = 1. 26) and is verysoluble in water. Under standard conditions 1 volume of water dissolvesabout 500 volumes of the gas. On warming such a solution the gasescapes, until at the boiling point the solution contains about 20% byweight of HCl. Further boiling will not drive out any more acid, but thesolution will distill with unchanged concentration. A more dilutesolution than this will lose water on boiling until it has reached thesame concentration, 20%, and will then distill unchanged. Under highpressure the gas can be liquefied, 28 atmospheres being required at 0°. Under these conditions it forms a colorless liquid which is not veryactive chemically. It boils at -80° and solidifies at -113°. Thesolution of the gas in water is used almost entirely in the place of thegas itself, since it is not only far more convenient but also moreactive. ~Chemical properties. ~ The most important chemical properties ofhydrochloric acid are the following: 1. _Action as an acid. _ In aqueous solution hydrochloric acid has verystrong acid properties; indeed, it is one of the strongest acids. Itacts upon oxides and hydroxides, converting them into salts: NaOH + HCl = NaCl + H_{2}O, CuO + 2HCl = CuCl_{2} + H_{2}O. It acts upon many metals, forming chlorides and liberating hydrogen: Zn + 2HCl = ZnCl_{2} + 2H, Al + 3HCl = AlCl_{3} + 3H. Unlike nitric and sulphuric acids it has no oxidizing action, so thatwhen it acts on metals hydrogen is always given off. 2. _Relation to combustion. _ Hydrochloric acid gas is not readilydecomposed, and is therefore neither combustible nor a supporter ofcombustion. 3. _Action on oxidizing agents. _ Although hydrochloric acid isincombustible, it can be oxidized under some circumstances, in whichcase the hydrogen combines with oxygen, while the chlorine is set free. Thus, when a solution of hydrochloric acid acts upon manganese dioxidepart of the chlorine is set free: MnO_{2} + 4HCl = MnCl_{2} + 2H_{2}O + 2Cl. ~Aqua regia. ~ It has been seen that when nitric acid acts as an oxidizingagent it usually decomposes, as represented in the equation 2HNO_{3} = H_{2}O + 2NO + 3O. The oxygen so set free may act on hydrochloric acid: 6HCl + 3O = 3H_{2}O + 6Cl. The complete equation therefore is 2HNO_{3} + 6HCl = 4H_{2}O + 2NO + 6Cl. When concentrated nitric and hydrochloric acids are mixed this reactiongoes on slowly, chlorine and some other substances not represented inthe equation being formed. The mixture is known as _aqua regia_ and iscommonly prepared by adding one volume of nitric acid to three volumesof hydrochloric acid. It acts more powerfully upon metals and othersubstances than either of the acids separately, and owes its strengthnot to acid properties but to the action of the nascent chlorine whichit liberates. Consequently, when it acts upon metals such as gold itconverts them into chlorides, and the reaction can be represented bysuch equations as Au + 3Cl = AuCl_{3}. ~Salts of hydrochloric acid, --chlorides. ~ The chlorides of all the metalsare known and many of them are very important compounds. Some of themare found in nature, and all can be prepared by the general method ofpreparing salts. Silver chloride, lead chloride, and mercurous chlorideare insoluble in water and acids, and can be prepared by addinghydrochloric acid to solutions of compounds of the respective elements. While the chlorides have formulas similar to the fluorides, theirproperties are often quite different. This is seen in the solubility ofthe salts. Those metals whose chlorides are insoluble form solublefluorides, while many of the metals which form soluble chlorides forminsoluble fluorides. ~Compounds of chlorine with oxygen and hydrogen. ~ Chlorine combines withoxygen and hydrogen to form four different acids. They are all quiteunstable, and most of them cannot be prepared in pure form; their saltscan easily be made, however, and some of them will be met with in thestudy of the metals. The formulas and names of these acids are asfollows: HClO hypochlorous acid. HClO_{2} chlorous acid. HClO_{3} chloric acid. HClO_{4} perchloric acid. ~Oxides of chlorine. ~ Two oxides are known, having the formulas Cl_{2}Oand ClO_{2}. They decompose very easily and are good oxidizing agents. BROMINE ~Historical. ~ Bromine was discovered in 1826 by the French chemistBallard, who isolated it from sea salt. He named it bromine (stench)because of its unbearable fumes. ~Occurrence. ~ Bromine occurs almost entirely in the form of bromides, especially as sodium bromide and magnesium bromide, which are found inmany salt springs and salt deposits. The Stassfurt deposits in Germanyand the salt waters of Ohio and Michigan are especially rich inbromides. ~Preparation of bromine. ~ The laboratory method of preparing bromine isessentially different from the commercial method. [Illustration Fig. 55] 1. _Laboratory method. _ As in the case of chlorine, bromine can beprepared by the action of hydrobromic acid (HBr) on manganese dioxide. Since hydrobromic acid is not an article of commerce, a mixture ofsulphuric acid and a bromide is commonly substituted for it. Thematerials are placed in a retort arranged as shown in Fig. 55. The endof the retort just touches the surface of the water in the test tube. Onheating, the bromine distills over and is collected in the coldreceiver. The equation is 2NaBr + 2H_{2}SO_{4} + MnO_{2} = Na_{2}SO_{4} + MnSO_{4} + 2H_{2}O + 2Br. 2. _Commercial method. _ Bromine is prepared commercially from the watersof salt wells which are especially rich in bromides. On passing acurrent of electricity through such waters the bromine is firstliberated. Any chlorine liberated, however, will assist in the reaction, since free chlorine decomposes bromides, as shown in the equation NaBr + Cl = NaCl + Br. When the water containing the bromine is heated, the liberated brominedistills over into the receiver. ~Physical properties. ~ Bromine is a dark red liquid about three times asheavy as water. Its vapor has a very offensive odor and is mostirritating to the eyes and throat. The liquid boils at 59° andsolidifies at -7°; but even at ordinary temperatures it evaporatesrapidly, forming a reddish-brown gas very similar to nitrogen peroxidein appearance. Bromine is somewhat soluble in water, 100 volumes ofwater under ordinary conditions dissolving 1 volume of the liquid. It isreadily soluble in carbon disulphide, forming a yellow solution. ~Chemical properties and uses. ~ In chemical action bromine is very similarto chlorine. It combines directly with many of the same elements withwhich chlorine unites, but with less energy. It combines with hydrogenand takes away the latter element from some of its compounds, but notso readily as does chlorine. Its bleaching properties are also lessmarked. Bromine finds many uses in the manufacture of organic drugs anddyestuffs and in the preparation of bromides. ~Hydrobromic acid (HBr). ~ When sulphuric acid acts upon a bromidehydrobromic acid is set free: 2NaBr + H_{2}SO_{4} = Na_{2}SO_{4} + 2HBr. At the same time some bromine is set free, as may be seen from the redfumes which appear, and from the odor. The explanation of this is foundin the fact that hydrobromic acid is much less stable than hydrochloricacid, and is therefore more easily oxidized. Concentrated sulphuric acidis a good oxidizing agent, and oxidizes a part of the hydrobromic acid, liberating bromine: H_{2}SO_{4} + 2HBr = 2H_{2}O + SO_{2} + 2Br. ~Preparation of pure hydrobromic acid. ~ A convenient way to make pure hydrobromic acid is by the action of bromine upon moist red phosphorus. This can be done with the apparatus shown in Fig. 56. Bromine is put into the dropping funnel A, and red phosphorus, together with enough water to cover it, is placed in the flask B. By means of the stopcock the bromine is allowed to flow drop by drop into the flask, the reaction taking place without the application of heat. The equations are (1) P + 3Br = PBr_{3}, (2) PBr_{3} + 3H_{2}O = P(OH)_{3} + 3HBr. [Illustration Fig. 56] The U-tube C contains glass beads which have been moistened with water and rubbed in red phosphorus. Any bromine escaping action in the flask acts upon the phosphorus in the U-tube. The hydrobromic acid is collected in the same way as hydrochloric acid. ~Properties. ~ Hydrobromic acid very strikingly resembles hydrochloric acidin physical and chemical properties. It is a colorless, strongly fuminggas, heavier than hydrochloric acid and, like it, is very soluble inwater. Under standard conditions 1 volume of water dissolves 610 volumesof the gas. Chemically, the chief point in which it differs fromhydrochloric acid is in the fact that it is much more easily oxidized, so that bromine is more readily set free from it than chlorine is fromhydrochloric acid. ~Salts of hydrobromic acid, --bromides. ~ The bromides are very similar tothe chlorides in their properties. Chlorine acts upon both bromides andfree hydrobromic acid, liberating bromine from them: KBr + Cl = KCl + Br, HBr + Cl = HCl + Br. Silver bromide is extensively used in photography, and the bromides ofsodium and potassium are used as drugs. ~Oxygen compounds. ~ No oxides of bromine are surely known, and bromine does not form so many oxygen acids as chlorine does. Salts of hypobromous acid (HBrO) and bromic acid (HBrO_{3}) are known. IODINE ~Historical. ~ Iodine was discovered in 1812 by Courtois in the ashes ofcertain sea plants. Its presence was revealed by its beautiful violetvapor, and this suggested the name iodine (from the Greek for violetappearance). ~Occurrence. ~ In the combined state iodine occurs in very small quantitiesin sea water, from which it is absorbed by certain sea plants, so thatit is found in their ashes. It occurs along with bromine in salt springsand beds, and is also found in Chili saltpeter. ~Preparation. ~ Iodine may be prepared in a number of ways, the principalmethods being the following: 1. _Laboratory method. _ Iodine can readily be prepared in the laboratoryfrom an iodide by the method used in preparing bromine, except thatsodium iodide is substituted for sodium bromide. It can also be made bypassing chlorine into a solution of an iodide. [Illustration: Fig. 57] 2. _Commercial method. _ Commercially iodine was formerly prepared fromseaweed (kelp), but is now obtained almost entirely from the deposits ofChili saltpeter. The crude saltpeter is dissolved in water and thesolution evaporated until the saltpeter crystallizes. The remainingliquors, known as the "mother liquors, " contain sodium iodate(NaIO_{3}), in which form the iodine is present in the saltpeter. Thechemical reaction by which the iodine is liberated from this compound isa complicated one, depending on the fact that sulphurous acid acts uponiodic acid, setting iodine free. This reaction is shown as follows: 2HIO_{3} + 5H_{2}SO_{3} = 5H_{2}SO_{4} + H_{2}O + 2I. ~Purification of iodine. ~ Iodine can be purified very conveniently in the following way. The crude iodine is placed in an evaporating dish E (Fig. 57), and the dish is set upon the sand bath S. The iodine is covered with the inverted funnel F, and the sand bath is gently heated with a Bunsen burner. As the dish becomes warm the iodine rapidly evaporates and condenses again on the cold surface of the funnel in shining crystals. This process, in which a solid is converted into a vapor and is again condensed into a solid without passing through the liquid state, is called _sublimation_. ~Physical properties. ~ Iodine is a purplish-black, shining, heavy solidwhich crystallizes in brilliant plates. Even at ordinary temperatures itgives off a beautiful violet vapor, which increases in amount as heat isapplied. It melts at 107° and boils at 175°. It is slightly soluble inwater, but readily dissolves in alcohol, forming a brown solution(tincture of iodine), and in carbon disulphide, forming a violetsolution. The element has a strong, unpleasant odor, though by no meansas irritating as that of chlorine and bromine. ~Chemical properties. ~ Chemically iodine is quite similar to chlorine andbromine, but is still less active than bromine. It combines directlywith many elements at ordinary temperatures. At elevated temperatures itcombines with hydrogen, but the reaction is reversible and the compoundformed is quite easily decomposed. Both chlorine and bromine displace itfrom its salts: KI + Br = KBr + I, KI + Cl = KCl + I. When even minute traces of iodine are added to thin starch paste a veryintense blue color develops, and this reaction forms a delicate test foriodine. Iodine is extensively used in medicine, especially in the formof a tincture. It is also largely used in the preparation of dyes andorganic drugs, iodoform, a substance used as an antiseptic, has theformula CHI_{3}. ~Hydriodic acid (HI). ~ This acid cannot be prepared in pure condition bythe action of sulphuric acid upon an iodide, since the hydriodic acidset free is oxidized by the sulphuric acid just as in the case ofhydrobromic acid, but to a much greater extent. It can be prepared inexactly the same way as hydrobromic acid, iodine being substituted forbromine. It can also be prepared by passing hydrosulphuric acid intowater in which iodine is suspended. The equation is H_{2}S + 2I = 2HI + S. The hydriodic acid formed in this way dissolves in the water. ~Properties and uses. ~ Hydriodic acid resembles the corresponding acids ofchlorine and bromine in physical properties, being a strongly fuming, colorless gas, readily soluble in water. Under standard conditions 1volume of water dissolves about 460 volumes of the gas. It is, however, more unstable than either hydrochloric or hydrobromic acids, and onexposure to the air it gradually decomposes in accordance with theequation 2HI + O = H_{2}O + 2I. Owing to the slight affinity between iodine and hydrogen the acid easilygives up its hydrogen and is therefore a strong reducing agent. This isseen in its action on sulphuric acid. The salts of hydriodic acid, the iodides, are, in general, similar tothe chlorides and bromides. Potassium iodide (KI) is the most familiarof the iodides and is largely used in medicine. ~Oxygen compounds. ~ Iodine has a much greater affinity for oxygen than has either chlorine or bromine. When heated with nitric acid it forms a stable oxide (I_{2}O_{5}). Salts of iodic acid (HIO_{3}) and periodic acid (HIO_{4}) are easily prepared, and the free acids are much more stable than the corresponding acids of the other members of this family. GAY-LUSSAC'S LAW OF VOLUMES In the discussion of the composition of hydrochloric acid it was statedthat one volume of hydrogen combines with one volume of chlorine to formtwo volumes of hydrochloric acid. With bromine and iodine similarcombining ratios hold good. These facts recall the simple volumerelations already noted in the study of the composition of steam andammonia. These relations may be represented graphically in the followingway: +---+ +----+ +------+ +------+ | H | + | Cl | = | H Cl | + | H Cl | +---+ +----+ +------+ +------+ +---+ +---+ +---+ +--------+ +--------+ | H | | H | + | O | = | H_{2}O | + | H_{2}O | +---+ +---+ +---+ +--------+ +--------+ +---+ +---+ +---+ +---+ +--------+ +--------+| H | | H | | H | + | N | = | NH_{3} | + | NH_{3} |+---+ +---+ +---+ +---+ +--------+ +--------+ In the early part of the past century Gay-Lussac, a distinguished Frenchchemist, studied the volume relations of many combining gases, andconcluded that similar relations always hold. His observations aresummed up in the following law: _When two gases combine chemically thereis always a simple ratio between their volumes, and between the volumeof either one of them and that of the product, provided it is a gas. _ Bya simple ratio is meant of course the ratio of small whole numbers, as1 : 2, 2 : 3. EXERCISES 1. How do we account for the fact that liquid hydrofluoric acid is notan electrolyte? 2. Why does sulphuric acid liberate hydrofluoric acid from its salts? 3. In the preparation of chlorine, what advantages are there in treatingmanganese dioxide with a mixture of sodium chloride and sulphuric acidrather than with hydrochloric acid? 4. Why must chlorine water be kept in the dark? 5. What is the derivation of the word nascent? 6. What substances studied are used as bleaching agents? To what is thebleaching action due in each case? 7. What substances studied are used as disinfecting agents? 8. What is meant by the statement that hydrochloric acid is one of thestrongest acids? 9. What is the meaning of the phrase _aqua regia_? 10. Cl_{2}O is the anhydride of what acid? 11. A solution of hydriodic acid on standing turns brown. How is thisaccounted for? 12. How can bromine vapor and nitrogen peroxide be distinguished fromeach other? 13. Write the equations for the reaction taking place when hydriodicacid is prepared from iodine, phosphorus, and water. 14. From their behavior toward sulphuric acid, to what class of agentsdo hydrobromic and hydriodic acids belong? 15. Give the derivation of the names of the elements of the chlorinefamily. 16. Write the names and formulas for the binary acids of the group inthe order of the stability of the acids. 17. What is formed when a metal dissolves in each of the following?nitric acid; dilute sulphuric acid; concentrated sulphuric acid;hydrochloric acid; aqua regia. 18. How could you distinguish between a chloride, a bromide, and aniodide? 19. What weight of sodium chloride is necessary to prepare sufficienthydrochloric acid to saturate 1 l. Of water under standard conditions? 20. On decomposition 100 l. Of hydrochloric acid would yield how manyliters of hydrogen and chlorine respectively, the gases being measuredunder the same conditions? Are your results in accord with theexperimental facts? CHAPTER XVII CARBON AND SOME OF ITS SIMPLER COMPOUNDS ~The family. ~ Carbon stands at the head of a family of elements in thefourth group in the periodic table. The resemblances between theelements of this family, while quite marked, are not so striking as inthe case of the elements of the chlorine family. With the exception ofcarbon, these elements are comparatively rare, and need not be taken upin detail in this chapter. Titanium will be referred to again inconnection with silicon which it very closely resembles. ~Occurrence. ~ Carbon is found in nature in the uncombined state in severalforms. The diamond is practically pure carbon, while graphite and coalare largely carbon, but contain small amounts of other substances. Itsnatural compounds are exceedingly numerous and occur as gases, liquids, and solids. Carbon dioxide is its most familiar gaseous compound. Natural gas and petroleum are largely compounds of carbon with hydrogen. The carbonates, especially calcium carbonate, constitute great strata ofrocks, and are found in almost every locality. All living organisms, both plant and animal, contain a large percentage of this element, andthe number of its compounds which go to make up all the vast variety ofanimate nature is almost limitless. Over one hundred thousand definitecompounds containing carbon have been prepared. In the free state carbonoccurs in three allotropic forms, two of which are crystalline and oneamorphous. ~Crystalline carbon. ~ Crystalline carbon occurs in two forms, --diamond andgraphite. 1. _Diamond. _ Diamonds are found in considerable quantities in severallocalities, especially in South Africa, the East Indies, and Brazil. Thecrystals belong to the regular system, but the natural stones do notshow this very clearly. When found they are usually covered with a roughcoating which is removed in the process of cutting. Diamond cutting iscarried on most extensively in Holland. The density of the diamond is 3. 5, and, though brittle, it is one of thehardest of substances. Black diamonds, as well as broken and imperfectstones which are valueless as gems, are used for grinding hardsubstances. Few chemical reagents have any action on the diamond, butwhen heated in oxygen or the air it blackens and burns, forming carbondioxide. Lavoisier first showed that carbon dioxide is formed by the combustionof the diamond; and Sir Humphry Davy in 1814 showed that this is theonly product of combustion, and that the diamond is pure carbon. ~The diamond as a gem. ~ The pure diamond is perfectly transparent and colorless, but many are tinted a variety of colors by traces of foreign substances. Usually the colorless ones are the most highly prized, although in some instances the color adds to the value; thus the famous Hope diamond is a beautiful blue. Light passing through a diamond is very much refracted, and to this fact the stone owes its brilliancy and sparkle. ~Artificial preparation of diamonds. ~ Many attempts have been made to produce diamonds artificially, but for a long time these always ended in failure, graphite and not diamonds being the product obtained. The French chemist Moissan, in his extended study of chemistry at high temperatures, finally succeeded (1893) in making some small ones. He accomplished this by dissolving carbon in boiling iron and plunging the crucible containing the mixture into water, as shown in Fig. 58. Under these conditions the carbon crystallized in the iron in the form of the diamond. The diamonds were then obtained by dissolving away the iron in hydrochloric acid. [Illustration: Fig. 58] 2. _Graphite. _ This form of carbon is found in large quantities, especially in Ceylon, Siberia, and in some localities of the UnitedStates and Canada. It is a shining black substance, very soft and greasyto the touch. Its density is about 2. 15. It varies somewhat inproperties according to the locality in which it is found, and is moreeasily attacked by reagents than is the diamond. It is also manufacturedby heating carbon with a small amount of iron (3%) in an electricfurnace. It is used in the manufacture of lead pencils and crucibles, asa lubricant, and as a protective covering for iron in the form of apolish or a paint. ~Amorphous carbon. ~ Although there are many varieties of amorphous carbonknown, they are not true allotropic modifications. They differ merely intheir degree of purity, their fineness of division, and in their mode ofpreparation. These substances are of the greatest importance, owing totheir many uses in the arts and industries. As they occur in nature, orare made artificially, they are nearly all impure carbon, the impuritydepending on the particular substance in question. 1. _Pure carbon. _ Pure amorphous carbon is best prepared by charringsugar. This is a substance consisting of carbon, hydrogen, and oxygen, the latter two elements being present in the ratio of one oxygen atom totwo of hydrogen. When sugar is strongly heated the oxygen and hydrogenare driven off in the form of water and pure carbon is left behind. Prepared in this way it is a soft, lustrous, very bulky, black powder. 2. _Coal and coke. _ Coals of various kinds were probably formed fromvast accumulations of vegetable matter in former ages, which becamecovered over with earthy material and were thus protected from rapiddecay. Under various natural agencies the organic matter was slowlychanged into coal. In anthracite these changes have gone the farthest, and this variety of coal is nearly pure carbon. Soft or bituminous coalscontain considerable organic matter besides carbon and mineralsubstances. When heated strongly out of contact with air the organicmatter is decomposed and the resulting volatile matter is driven off inthe form of gases and vapors, and only the mineral matter and carbonremain behind. The gaseous product is chiefly illuminating gas and thesolid residue is _coke_. Some of the coke is found as a dense cake onthe sides and roof of the retort. This is called retort carbon and isquite pure. 3. _Charcoal. _ This is prepared from wood in the same way that coke ismade from coal. When the process is carried on in retorts the productsexpelled by the heat are saved. Among these are many valuable substancessuch as wood alcohol and acetic acid. Where timber is abundant theprocess is carried out in a wasteful way, by merely covering piles ofwood with sod and setting the wood on fire. Some wood burns and the heatfrom this decomposes the wood not burned, forming charcoal from it. Thecharcoal, of course, contains the mineral part of the wood from which itis formed. 4. _Bone black. _ This is sometimes called animal charcoal, and is madeby charring bones and animal refuse. The organic part of the materialsis thus decomposed and carbon is left in a very finely divided state, scattered through the mineral part which consists largely of calciumphosphate. For some uses this mineral part is removed by treatment withhydrochloric acid and prolonged washing. 5. _Lampblack. _ Lampblack and soot are products of imperfect combustionof oil and coal, and are deposited from a smoky flame on a cold surface. The carbon in this form is very finely divided and usually containsvarious oily materials. ~Properties. ~ While the various forms of carbon differ in many properties, especially in color and hardness, yet they are all odorless, tastelesssolids, insoluble in water and characterized by their stability towardsheat. Only in the intense heat of the electric arc does carbonvolatilize, passing directly from the solid state into a vapor. Owing tothis fact the inside surface of an incandescent light bulb after beingused for some time becomes coated with a dark film of carbon. It is notacted on at ordinary temperatures by most reagents, but at a highertemperature it combines directly with many of the elements, formingcompounds called _carbides_. When heated in the presence of sufficientoxygen it burns, forming carbon dioxide. ~Uses of carbon. ~ The chief use of amorphous carbon is for fuel to furnishheat and power for all the uses of civilization. An enormous quantity ofcarbon in the form of the purer coals, coke, and charcoal is used as areducing agent in the manufacture of the various metals, especially inthe metallurgy of iron. Most of the metals are found in nature asoxides, or in forms which can readily be converted into oxides. Whenthese oxides are heated with carbon the oxygen is abstracted, leavingthe metal. Retort carbon and coke are used to make electric lightcarbons and battery plates, while lampblack is used for indelible inks, printer's ink, and black varnishes. Bone black and charcoal have theproperty of absorbing large volumes of certain gases, as well as smalleramounts of organic matter; hence they are used in filters to removenoxious gases and objectionable colors and odors from water. Bone blackis used extensively in the sugar refineries to remove coloring matterfrom the impure sugars. ~Chemistry of carbon compounds. ~ Carbon is remarkable for the very largenumber of compounds which it forms with the other elements, especiallywith oxygen and hydrogen. Compounds containing carbon are more numerousthan all others put together, and the chemistry of these substancespresents peculiarities not met with in the study of other substances. For these reasons the systematic study of carbon compounds, or of_organic chemistry_ as it is usually called, must be deferred until thestudent has gained some knowledge of the chemistry of other elements. Anacquaintance with a few of the most familiar carbon compounds is, however, essential for the understanding of the general principles ofchemistry. ~Compounds of carbon with hydrogen, --the hydrocarbons. ~ Carbon unites withhydrogen to form a very large number of compounds called _hydrocarbons_. Petroleum and natural gas are essentially mixtures of a great variety ofthese hydrocarbons. Many others are found in living plants, and stillothers are produced by the decay of organic matter in the absence ofair. Only two of them, methane and acetylene, will be discussed here. ~Methane~ (_marsh gas_) (CH_{4}). This is one of the most important ofthese hydrocarbons, and constitutes about nine tenths of natural gas. Asits name suggests, it is formed in marshes by the decay of vegetablematter under water, and bubbles of the gas are often seen to rise whenthe dead leaves on the bottom of pools are stirred. It also collects inmines, and, when mixed with air, is called _fire damp_ by the minersbecause of its great inflammability, damp being an old name for a gas. It is formed when organic matter, such as coal or wood, is heated inclosed vessels, and is therefore a principal constituent of coal gas. ~Preparation. ~ Methane is prepared in the laboratory by heating sodium orcalcium acetate with soda-lime. Equal weights of fused sodium acetateand soda-lime are thoroughly dried, then mixed and placed in agood-sized, hard-glass test tube fitted with a one-holed stopper anddelivery tube. The mixture is gradually heated, and when the air hasbeen displaced from the tube the gas is collected in bottles bydisplacement of water. Soda-lime is a mixture of sodium and calciumhydroxides. Regarding it as sodium hydroxide alone, the equation is NaC_{2}H_{3}O_{2} + NaOH = Na_{2}CO_{3} + CH_{4}. ~Properties. ~ Methane is a colorless, odorless gas whose density is 0. 55. It is difficult to liquefy, boiling at -155° under standard pressure, and is almost insoluble in water. It burns with a pale blue flame, liberating much heat, and when mixed with oxygen is very explosive. ~Davy's safety lamp. ~ In 1815 Sir Humphry Davy invented a lamp for the useof miners, to prevent the dreadful mine explosions then common, due tomethane mixed with air. The invention consisted in surrounding the upperpart of the common miner's lamp with a mantle of wire gauze and thelower part with glass (Fig. 59). It has been seen that two gases willnot combine until raised to their kindling temperature, and if whilecombining they are cooled below this point, the combination ceases. Aflame will not pass through a wire gauze because the metal, being a goodconductor of heat, takes away so much heat from the flame that the gasesare cooled below the kindling temperature. When a lamp so protected isbrought into an explosive mixture the gases inside the wire mantle burnin a series of little explosions, giving warning to the miner that theair is unsafe. [Illustration: Fig. 59] ~Acetylene~ (C_{2}H_{2}). This is a colorless gas usually having adisagreeable odor due to impurities. It is now made in large quantitiesfrom calcium carbide (CaC_{2}). This substance is formed when coal andlime are heated together in an electric furnace. When treated with waterthe carbide is decomposed, yielding acetylene: CaC_{2} + 2H_{2}O = C_{2}H_{2} + Ca(OH)_{2}. Under ordinary conditions the gas burns with a very smoky flame; inburners constructed so as to secure a large amount of oxygen it burnswith a very brilliant white light, and hence is used as an illuminant. ~Laboratory preparation. ~ The gas can be prepared readily in a generatorsuch as is shown in Fig. 60. The inner tube contains fragments ofcalcium carbide, while the outer one is filled with water. As long asthe stopcock is closed the water cannot rise in the inner tube. When thestopcock is open the water rises, and, coming into contact with thecarbide in the inner tube, generates acetylene. This escapes through thestopcock, and after the air has been expelled may be lighted as itissues from the burner. [Illustration: Fig. 60] Carbon forms two oxides, namely, carbon dioxide (CO_{2}) and carbonmonoxide (CO). ~Carbon dioxide~ (CO_{2}). Carbon dioxide is present in the air to theextent of about 3 parts in 10, 000, and this apparently small amount isof fundamental importance in nature. In some localities it escapes fromthe earth in great quantities, and many spring waters carry largeamounts of it in solution. When these highly charged spring waters reachthe surface of the earth, and the pressure on them is removed, thecarbon dioxide escapes with effervescence. It is a product of theoxidation of all organic matter, and is therefore formed in fires aswell as in the process of decay. It is thrown off from the lungs of allanimals in respiration, and is a product of many fermentation processessuch as vinegar making and brewing. Combined with metallic oxides itforms vast deposits of carbonates in nature. ~Preparation. ~ In the laboratory carbon dioxide is always prepared by theaction of an acid upon a carbonate, usually calcium carbonate, theapparatus shown in Fig. 39 serving the purpose very well. This reactionmight be expected to produce carbonic acid, thus: CaCO_{3} + 2HCl = CaCl_{2} + H_{2}CO_{3}. Carbonic acid is very unstable, however, and decomposes into itsanhydride, CO_{2}, and water, thus: H_{2}CO_{3} = H_{2}O + CO_{2}. The complete reaction is represented by the equation CaCO_{3} + 2HCl = CaCl_{2} + CO_{2} + H_{2}O. ~Physical properties. ~ Carbon dioxide is a colorless, practically odorlessgas whose density is 1. 5. Its weight may be inferred from the fact thatit can be siphoned, or poured like water, from one vessel downward intoanother. At 15° and under ordinary pressure it dissolves in its ownvolume of water and imparts a somewhat biting, pungent taste to it. Itis easily condensed, and is now prepared commercially in this form bypumping the gas into steel cylinders (see Fig. 6) which are kept coldduring the process. When the liquid is permitted to escape into the airpart of it instantly evaporates, and in so doing absorbs so much heatthat another portion is solidified, the solid form strikingly resemblingsnow in appearance. This snow is very cold and mercury can easily befrozen with it. ~Solid carbon dioxide. ~ Cylinders of liquid carbon dioxide areinexpensive, and should be available in every school. To demonstrate theproperties of solid carbon dioxide, the cylinder should be placed acrossthe table and supported in such a way that the stopcock end is severalinches lower than the other end. A loose bag is made by holding thecorners of a handkerchief around the neck of the stopcock, and the cockis then turned on so that the gas rushes out in large quantities. Veryquickly a considerable quantity of the snow collects in thehandkerchief. To freeze mercury, press a piece of filter paper into asmall evaporating dish and pour the mercury upon it. Coil a flat spiralupon the end of a wire, and dip the spiral into the mercury. Place aquantity of solid carbon dioxide upon the mercury and pour 10 cc. -15 cc. Of ether over it. In a minute or two the mercury will solidify and maybe removed from the dish by the wire serving as a handle. The filterpaper is to prevent the mercury from sticking to the dish; the etherdissolves the solid carbon dioxide and promotes its rapid conversioninto gas. ~Chemical properties. ~ Carbon dioxide is incombustible, since it is, likewater, a product of combustion. It does not support combustion, as doesnitrogen peroxide, because the oxygen in it is held in very firmchemical union with the carbon. Very strong reducing agents, such ashighly heated carbon, can take away half of its oxygen: CO_{2} + C = 2CO. ~Uses. ~ The relation of carbon dioxide to plant life has been discussed ina previous chapter. Water highly charged with carbon dioxide is used formaking soda water and similar beverages. Since it is a non-supporter ofcombustion and can be generated readily, carbon dioxide is also used asa fire extinguisher. Some of the portable fire extinguishers are simplydevices for generating large amounts of the gas. It is not necessarythat all the oxygen should be kept away from the fire in order tosmother it. A burning candle is extinguished in air which contains only2. 5% of carbon dioxide. ~Carbonic acid~ (H_{2}CO_{3}). Like most of the oxides of the non-metallicelements, carbon dioxide is an acid anhydride. It combines with water toform an acid of the formula H_{2}CO_{3}, called carbonic acid: H_{2}O + CO_{2} = H_{2}CO_{3}. The acid is, however, very unstable and cannot be isolated. Only a verysmall amount of it is actually formed when carbon dioxide is passed intowater, as is evident from the small solubility of the gas. If, however, a base is present in the water, salts of carbonic acid are formed, andthese are quite stable: 2NaOH + H_{2}O + CO_{2} = Na_{2}CO_{3} + 2H_{2}O. ~Action of carbon dioxide on bases. ~ This conduct is explained by theprinciples of reversible reactions. The equation H_{2}O +CO_{2} H_{2}CO_{3} is a reversible equation, and the extent to which the reactionprogresses depends upon the relative concentrations of each of the threefactors in it. Equilibrium is ordinarily reached when very littleH_{2}CO_{3} is formed. If a base is present in the water to combine withthe H_{2}CO_{3} as fast as it is formed, all of the CO_{2} is convertedinto H_{2}CO_{3}, and thence into a carbonate. ~Salts of carbonic acid, --carbonates. ~ The carbonates form a veryimportant class of salts. They are found in large quantities in nature, and are often used in chemical processes. Only the carbonates of sodium, potassium, and ammonium are soluble, and these can be made by the actionof carbon dioxide on solutions of the bases, as has just been explained. The insoluble carbonates are formed as precipitates when soluble saltsare treated with a solution of a soluble carbonate. Thus the insolublecalcium carbonate can be made by bringing together solutions of calciumchloride and sodium carbonate: CaCl_{2} + Na_{2}CO_{3} = CaCO_{3} + 2NaCl. Most of the carbonates are decomposed by heat, yielding an oxide of themetal and carbon dioxide. Thus lime (calcium oxide) is made by stronglyheating calcium carbonate: CaCO_{3} = CaO + CO_{2}. ~Acid carbonates. ~ Like all acids containing two acid hydrogen atoms, carbonic acid can form both normal and acid salts. The acid carbonatesare made by treating a normal carbonate with an excess of carbonic acid. With few exceptions they are very unstable, heat decomposing them evenwhen in solution. ~Action of carbon dioxide on calcium hydroxide. ~ If carbon dioxide ispassed into clear lime water, calcium carbonate is at firstprecipitated: H_{2}O + CO_{2} = H_{2}CO_{3}, Ca(OH)_{2} + H_{2}CO_{3} = CaCO_{3} + 2H_{2}O. Advantage is taken of this reaction in testing for the presence ofcarbon dioxide, as already explained in the chapter on the atmosphere. If the current of carbon dioxide is continued, the precipitate soondissolves, because the excess of carbonic acid forms calcium acidcarbonate which is soluble: CaCO_{3} + H_{2}CO_{3} = Ca(HCO_{3})_{2}. If now the solution is heated, the acid carbonate is decomposed andcalcium carbonate once more precipitated: Ca(HCO_{3})_{2} = CaCO_{3} + H_{2}CO_{3}. ~Carbon monoxide (CO). ~ Carbon monoxide can be made in a number of ways, the most important of which are the three following: 1. _By the partial oxidation of carbon. _ If a slow current of air isconducted over highly heated carbon, the monoxide is formed, thus: C + O = CO It is therefore often formed in stoves when the air draught isinsufficient. Water gas, which contains large amounts of carbonmonoxide, is made by partially oxidizing carbon with steam: C + H_{2}O = CO + 2H. 2. _By the partial reduction of carbon dioxide. _ When carbon dioxide isconducted over highly heated carbon it is reduced to carbon monoxide bythe excess of carbon: CO_{2} + C = 2CO. When coal is burning in a stove or grate carbon dioxide is at firstformed in the free supply of air, but as the hot gas rises through theglowing coal it is reduced to carbon monoxide. When the carbon monoxidereaches the free air above the coal it takes up oxygen to form carbondioxide, burning with the blue flame so familiar above a bed of coals, especially in the case of hard coals. 3. _By the decomposition of oxalic acid. _ In the laboratory carbonmonoxide is usually prepared by the action of concentrated sulphuricacid upon oxalic acid. The latter substance has the formulaC_{2}H_{2}O_{4}. The sulphuric acid, owing to its affinity for water, decomposes the oxalic acid, as represented in the equation C_{2}H_{2}O_{4} + (H_{2}SO_{4}) = (H_{2}SO_{4}) + H_{2}O + CO_{2} + CO. ~Properties. ~ Carbon monoxide is a light, colorless, almost odorless gas, very difficult to liquefy. Chemically it is very active, combiningdirectly with a great many substances. It has a great affinity foroxygen and is therefore combustible and a good reducing agent. Thus, ifcarbon monoxide is passed over hot copper oxide, the copper is reducedto the metallic state: CuO + CO = Cu + CO_{2}. When inhaled it combines with the red coloring matter of the blood andin this way prevents the absorption of oxygen, so that even a smallquantity of the gas may prove fatal. [Illustration: Fig. 61] ~The reducing power of carbon monoxide. ~ Fig. 61 illustrates a method of showing the reducing power of carbon monoxide. The gas is generated by gently heating 7 or 8 g. Of oxalic acid with 25 cc. Of concentrated sulphuric acid in a 200 cc. Flask A. The bottle B contains a solution of sodium hydroxide, which removes the carbon dioxide formed along with the monoxide. C contains a solution of calcium hydroxide to show that the carbon dioxide is completely removed. E is a hard-glass tube containing 1 or 2 g. Of copper oxide, which is heated by a burner. The black copper oxide is reduced to reddish metallic copper by the carbon monoxide, which is thereby changed to carbon dioxide. The presence of the carbon dioxide is shown by the precipitate in the calcium hydroxide solution in D. Any unchanged carbon monoxide is collected over water in F. ~Carbon disulphide~ (CS_{2}). Just as carbon combines with oxygen to formcarbon dioxide, so it combines with sulphur to form carbon disulphide(CS_{2}). This compound has been described in the chapter on sulphur. ~Hydrocyanic acid~ (_prussic acid_)(HCN). Under the proper conditionscarbon unites with nitrogen and hydrogen to form the acid HCN, calledhydrocyanic acid. It is a weak, volatile acid, and is therefore easilyprepared by treating its salts with sulphuric acid: KCN + H_{2}SO_{4} = KHSO_{4} + HCN. It is most familiar as a gas, though it condenses to a colorless liquidboiling at 26°. It has a peculiar odor, suggesting bitter almonds, andis extremely poisonous either when inhaled or when taken into thestomach. A single drop may cause death. It dissolves readily in water, its solution being commonly called prussic acid. The salts of hydrocyanic acid are called _cyanides_, the cyanides ofsodium and potassium being the best known. These are white solids andare extremely poisonous. ~Solutions of potassium cyanide are alkaline. ~ A solution of potassiumcyanide turns red litmus blue, and must therefore contain hydroxyl ions. The presence of these ions is accounted for in the following way. Although water is so little dissociated into its ions H^{+} and OH^{-}that for most purposes we may neglect the dissociation, it isnevertheless measurably dissociated. Hydrocyanic acid is one of theweakest of acids, and dissociates to an extremely slight extent. When acyanide such as potassium cyanide dissolves it freely dissociates, andthe CN^{-} ions must come to an equilibrium with the H^{+} ions derivedfrom the water: H^{+} + CN^{-} HCN. The result of this equilibrium is that quite a number of H^{+} ions fromthe water are converted into undissociated HCN molecules. But for everyH^{+} ion so removed an OH^{-} ion remains free, and this will give thesolution alkaline properties. EXERCISES 1. How can you prove that the composition of the different allotropicforms of carbon is the same? 2. Are lampblack and bone black allotropic forms of carbon? Will equalamounts of heat be liberated in the combustion of 1 g. Of each? 3. How could you judge of the relative purity of different forms ofcarbon? 4. Apart from its color, why should carbon be useful in the preparationof inks and paints? 5. Could asbestos fibers be used to replace the wire in a safety lamp? 6. Why do most acids decompose carbonates? 7. What effect would doubling the pressure have upon the solubility ofcarbon dioxide in water? 8. What compound would be formed by passing carbon dioxide into asolution of ammonium hydroxide? Write the equation. 9. Write equations for the preparation of K_{2}CO_{3}; of BaCO_{3}; ofMgCO_{3}. 10. In what respects are carbonic and sulphurous acids similar? 11. Give three reasons why the reaction which takes place when asolution of calcium acid carbonate is heated, completes itself. 12. How could you distinguish between carbonates and sulphites? 13. How could you distinguish between oxygen, hydrogen, nitrogen, nitrous oxide, and carbon dioxide? 14. Could a solution of sodium hydroxide be substituted for the solutionof calcium hydroxide in testing for carbon dioxide? 15. What weight of sodium hydroxide is necessary to neutralize thecarbonic acid formed by the action of hydrochloric acid on 100 g. Ofcalcium carbonate? 16. What weight of calcium carbonate would be necessary to preparesufficient carbon dioxide to saturate 10 l. Of water at 15° and underordinary pressure? 17. On the supposition that calcium carbide costs 12 cents a kilogram, what would be the cost of an amount sufficient to generate 100 l. Ofacetylene measured at 20° and 740 mm. ? 18. How would the volume of a definite amount of carbon monoxide comparewith the volume of carbon dioxide formed by its combustion, themeasurements being made under the same conditions? CHAPTER XVIII FLAMES, --ILLUMINANTS ~Conditions necessary for flames. ~ It has been seen that when twosubstances unite chemically, with the production of light and heat, theact of union is called combustion. When one of the substances undergoingcombustion remains solid at the temperature occasioned by thecombustion, light may be given off, but there is no flame. Thus ironwire burning in oxygen throws off a shower of sparks and is brilliantlyincandescent, but no flame is seen. When, however, both of thesubstances are gases or vapors at the temperature reached in thecombustion, the act of union is accompanied by a flame. ~Flames from burning liquids or solids. ~ Many substances which are liquidsor solids at ordinary temperatures burn with a flame because the heat ofcombustion vaporizes them slowly, and the flame is due to the union ofthis vapor with the gas supporting the combustion. ~Supporter of combustion. ~ That gas which surrounds the flame andconstitutes the atmosphere in which the combustion occurs is said tosupport the combustion. The other gas which issues into this atmosphereis said to be the combustible gas. Thus, in the ordinary combustion ofcoal gas in the air the coal gas is said to be combustible, while theair is regarded as the supporter of combustion. These terms are entirelyrelative, however, for a jet of air issuing into an atmosphere of coalgas will burn when ignited, the coal gas supporting the combustion. Ordinarily, when we say that a gas is combustible we mean that it iscombustible in an atmosphere of air. [Illustration: Fig. 62] ~Either gas may be the supporter of combustion. ~ That the terms_combustible_ and _supporter of combustion_ are merely relative may beshown in the following way: A lamp chimney A is fitted with a cork andglass tubes, as shown in Fig. 62. The tube C should have a diameter offrom 12 to 15 mm. A thin sheet of asbestos in which is cut a circularopening about 2 cm. In diameter is placed over the top of the chimney. The opening in the asbestos is closed with the palm of the hand, and gasis admitted to the chimney through the tube B. The air in the chimneyis soon expelled through the tube C, and the gas itself is thenlighted at the lower end of this tube. The hand is now removed from theopening in the asbestos, when the flame at the end of the tube at oncerises and appears at the end within the chimney, as shown in the figure. The excess of coal gas now escapes from the opening in the asbestos andmay be lighted. The flame at the top of the asbestos board is due to thecombustion of coal gas in air, while the flame within the chimney is dueto the combustion of air in coal gas, the air being drawn up through thetube by the escaping gas. ~Appearance of flames. ~ The flame caused by the union of hydrogen andoxygen is almost colorless and invisible. Chlorine and hydrogen combinewith a pale violet flame, carbon monoxide burns in oxygen with a blueflame, while ammonia burns with a deep yellow flame. The color andappearance of flames are therefore often quite characteristic of theparticular combustion which occasions them. ~Structure of flames. ~ When the gas undergoing combustion issues from around opening into an atmosphere of the gas supporting combustion, as isthe case with the burning Bunsen burner (Fig. 63), the flame isgenerally conical in outline. It consists of several distinct cones, one within the other, the boundary between them being marked bydifferences of color or luminosity. In the simplest flame, of whichhydrogen burning in oxygen is a good example, these cones are two innumber, --an inner one, formed by unburned gas, and an outer one, usuallymore or less luminous, consisting of the combining gases. This outer oneis in turn surrounded by a third envelope of the products of combustion;this envelope is sometimes invisible, as in the present case, but issometimes faintly luminous. The lower part of the inner cone of theflame is quite cool and consists of unburned gas. Toward the top of theinner cone the gas has become heated to a high temperature by theburning envelope surrounding it. On reaching the supporter of combustionon the outside it is far above its kindling temperature, and combustionfollows with the evolution of much heat. The region of combustion justoutside the inner cone is therefore the hottest part of the flame. [Illustration: Fig. 63] ~Oxidizing and reducing flames. ~ Since the tip of the outside coneconsists of very hot products of combustion mixed with oxygen from theair, a substance capable of oxidation placed in this part of the flamebecomes very hot and is easily oxidized. The oxygen with which itcombines comes, of course, from the atmosphere, and not from theproducts of combustion. This outer tip of the flame is called the_oxidizing flame_. At the tip of the inner cone the conditions are quite different. Thisregion consists of a highly heated combustible gas, which has not yetreached a supply of oxygen. If a substance rich in oxygen, such as a metallic oxide, is placed inthis region of the flame, the heated gases combine with its oxygen andthe substance is reduced. This part of the flame is called the _reducingflame_. These flames are used in testing certain substances, especiallyminerals. For this purpose they are produced by blowing into a smallluminous Bunsen flame from one side through a blowpipe. This is a tubeof the shape shown in Fig. 64. The flame is directed in any desired wayand has the oxidizing and reducing regions very clearly marked (Fig. 65). It is non-luminous from the same causes which render the openBunsen burner flame non-luminous, the gases from the lungs serving tofurnish oxygen and to dilute the combustible gas. [Illustration: Fig. 64] [Illustration: Fig. 65] ~Luminosity of flames. ~ The luminosity of flames is due to a number ofdistinct causes, and may therefore be increased or diminished in severalways. 1. _Presence of solid matter. _ The most obvious of these causes is thepresence in the flame of incandescent solid matter. Thus chalk dustsifted into a non-luminous flame renders it luminous. When hydrocarbonsform a part of the combustible gas, as they do in nearly allilluminating gases and oils, some carbon is usually set free in theprocess of combustion. This is made very hot by the flame and becomesincandescent, giving out light. In a well-regulated flame it isafterward burned up, but when the supply of oxygen is insufficient itescapes from the flame as lampblack or soot. That it is temporarilypresent in a well-burning luminous flame may be demonstrated by holdinga cold object, such as a small evaporating dish, in the flame for a fewseconds. This cold object cools the carbon below its kindlingtemperature, and it is deposited on the object as soot. 2. _Pressure. _ A second factor in the luminosity of flames is thepressure under which the gases are burning. Under increased pressurethere is more matter in a given volume of a gas, and the chemical actionis more energetic than when the gases are rarefied. Consequently thereis more heat and light. A candle burning on a high mountain gives lesslight than when it burns at the sea level. If the gas is diluted with a non-combustible gas, the effect is the sameas if it is rarefied, for under these conditions there is lesscombustible gas in a given volume. 3. _Temperature. _ The luminosity also depends upon the temperatureattained in the combustion. In general the hotter the flame the greaterthe luminosity; hence cooling the gases before combustion diminishes theluminosity of the flame they will make, because it diminishes thetemperature attained in the combustion. Thus the luminosity of theBunsen flame is largely diminished by the air drawn up with the gas. This is due in part to the fact that the burning gas is diluted andcooled by the air drawn in. The oxygen thus introduced into the flamealso causes the combustion of the hot particles of carbon which wouldotherwise tend to make the flame luminous. ~Illuminating and fuel gases. ~ A number of mixtures of combustible gases, consisting largely of carbon compounds and hydrogen, find extensive usefor the production of light and heat. The three chief varieties are coalgas, water gas, and natural gas. The use of acetylene gas has alreadybeen referred to. ~Coal gas. ~ Coal gas is made by heating bituminous coal in large retortsout of contact with the air. Soft or bituminous coal contains, inaddition to large amounts of carbon, considerable quantities ofcompounds of hydrogen, oxygen, nitrogen, and sulphur. When distilled thenitrogen is liberated partly in the form of ammonia and cyanides andpartly as free nitrogen gas; the sulphur is converted into hydrogensulphide, carbon disulphide, and oxides of sulphur; the oxygen intowater and oxides of carbon. The remaining hydrogen is set free partlyas hydrogen and partly in combination with carbon in the form ofhydrocarbons. The most important of these is methane, with smallerquantities of many others, some of which are liquids or solids atordinary temperatures. The great bulk of the carbon remains behind ascoke and retort carbon. ~The manufacture of coal gas. ~ In the manufacture of coal gas it isnecessary to separate from the volatile constituents formed by theheating of the coal all those substances which are either solid orliquid at ordinary temperature, since these would clog the gas pipes. Certain gaseous constituents, such as hydrogen sulphide and ammonia, must also be removed. The method used to accomplish this is shown inFig. 66. The coal is heated in air-tight retorts illustrated by A. Thevolatile products escape through the pipe X and bubble into the tarryliquid in the large pipe B, known as the _hydraulic main_, which runsat right angles to the retorts. Here is deposited the greater portion ofthe solid and liquid products, forming a tarry mass known as _coal tar_. Much of the ammonia also remains dissolved in this liquid. The partiallypurified gas then passes into the pipes C, which serve to cool it andfurther remove the solid and liquid matter. The gas then passes intoD, which is filled with coke over which a jet of water is sprayed. Thewater still further cools the gas and at the same time partially removessuch gaseous products as hydrogen sulphide and ammonia, which aresoluble in water. In E the gas passes over some material such as lime, which removes the last portions of the sulphur compounds as well as muchof the carbon dioxide present. From E the gas passes into the largegas holder F, from which it is distributed through pipes to the placeswhere it is burned. [Illustration: Fig. 66] One ton of good gas coal yields approximately 10, 000 cu. Ft. Of gas, 1400 lb. Of coke, 120 lb. Of tar, and 20 gal. Of ammoniacal liquor. Not only is the ammonia obtained in the manufacture of the gas of great importance, but the coal tar also serves as the source of many very useful substances, as will be explained in Chapter XXXII. ~Water gas. ~ Water gas is essentially a mixture of carbon monoxide andhydrogen. It is made by passing steam over very hot anthracite coal, when the reaction shown in the following equation takes place: C + H_{2}O = CO + 2H. When required merely to produce heat the gas is at once ready for use. When made for illuminating purposes it must be enriched, that is, illuminants must be added, since both carbon monoxide and hydrogen burnwith non-luminous flames. This is accomplished by passing it intoheaters containing highly heated petroleum oils. The gas takes uphydrocarbon gases formed in the decomposition of the petroleum oils, which make it burn with a luminous flame. Water gas is very effective as a fuel, since both carbon monoxide andhydrogen burn with very hot flames. It has little odor and is verypoisonous. Its use is therefore attended with some risk, since leaks inpipes are very likely to escape notice. ~Natural gas. ~ This substance, so abundant in many localities, varies muchin composition, but is composed principally of methane. When used forlighting purposes it is usually burned in a burner resembling an openBunsen, the illumination being furnished by an incandescent mantle. Thisis the case in the familiar Welsbach burner. Contrary to statementsfrequently made, natural gas contains no free hydrogen. TABLE SHOWING COMPOSITION OF GASES =====================+================+========+========+========== | PENNSYLVANIA | COAL | WATER | ENRICHED | NATURAL | GAS | GAS | WATER | GAS | | | GAS---------------------+----------------+--------+--------+----------Hydrogen | | 41. 3 | 52. 88 | 30. 00Methane | 90. 64 | 43. 6 | 2. 16 | 24. 00Illuminants | | 3. 9 | | 12. 05Carbon monoxide | | 6. 4 | 36. 80 | 29. 00Carbon dioxide | 0. 30 | 2. 0 | 3. 47 | 0. 30Nitrogen | 9. 06 | 1. 2 | 4. 69 | 2. 50Oxygen | | 0. 3 | | 1. 50Hydrocarbon vapors | | 1. 5 | | 1. 50=====================+================+========+========+========== These are analyses of actual samples, and may be taken as about the average for the various kinds of gases. Any one of these may vary considerably. The nitrogen and oxygen in most cases is due to a slight admixture of air which is difficult to exclude entirely in the manufacture and handling of gases. ~Fuels. ~ A variety of substances are used as fuels, the most important ofthem being wood, coal, and the various gases mentioned above. Woodconsists mainly of compounds of carbon, hydrogen, and oxygen. Thecomposition of coal and the fuel gases has been given. Since these fuelsare composed principally of carbon and hydrogen or their compounds, thechief products of combustion are carbon dioxide and water. The practiceof heating rooms with portable gas or oil stoves with no provision forremoving the products of combustion is to be condemned, since the carbondioxide is generated in sufficient quantities to render the air unfitfor breathing. Rooms so heated also become very damp from the largeamount of water vapor formed in the combustion, and which in coldweather condenses on the window glass, causing the glass to "sweat. "Both coal and wood contain a certain amount of mineral substances whichconstitute the ashes. ~The electric furnace. ~ In recent years electric furnaces have come intowide use in operations requiring a very high temperature. Temperaturesas high as 3500° can be easily reached, whereas the hottest oxyhydrogenflame is not much above 2000°. These furnaces are constructed on one oftwo general principles. [Illustration: Fig. 67] 1. _Arc furnaces. _ In the one type the source of heat is an electric arcformed between carbon electrodes separated a little from each other, asshown in Fig. 67. The substance to be heated is placed in a vessel, usually a graphite crucible, just below the arc. The electrodes andcrucible are surrounded by materials which fuse with great difficulty, such as magnesium oxide, the walls of the furnace being so shaped as toreflect the heat downwards upon the contents of the crucible. [Illustration: Fig. 68] 2. _Resistance furnaces. _ In the other type of furnace the heat isgenerated by the resistance offered to the current in its passagethrough the furnace. In its simplest form it may be represented by Fig. 68. The furnace is merely a rectangular box built up of loose bricks. The electrodes E, each consisting of a bundle of carbon rods, areintroduced through the sides of the furnace. The materials to be heated, C, are filled into the furnace up to the electrodes, and a layer ofbroken coke is arranged so as to extend from one electrode to the other. More of the charge is then placed on top of the coke. In passing throughthe broken coke the electrical current encounters great resistance. Thisgenerates great heat, and the charge surrounding the coke is brought toa very high temperature. The advantage of this type of furnace is thatthe temperature can be regulated to any desired intensity. EXERCISES 1. Why does charcoal usually burn with no flame? How do you account forthe flame sometimes observed when it burns? 2. How do you account for the fact that a candle burns with a flame? 3. What two properties must the mantle used in the Welsbach lamppossess? 4. (a) In what respects does the use of the Welsbach mantle resemblethat of lime in the calcium light? (b) If the mantle were made ofcarbon, would it serve the same purpose? 5. Would anthracite coal be suitable for the manufacture of coal gas? 6. How could you prove the formation of carbon dioxide and water in thecombustion of illuminating gases? 7. Suggest a probable way in which natural gas has been formed. 8. Coal frequently contains a sulphide of iron. (a) What two sulphurcompounds are likely to be formed when gas is made from such coal? (b)Suggest some suitable method for the removal of these compounds. 9. Why does the use of the bellows on the blacksmith's forge cause amore intense heat? 10. What volume of oxygen is necessary to burn 100 l. Of marsh gas andwhat volume of carbon dioxide would be formed, all of the gases beingmeasured under standard conditions? 11. Suppose a cubic meter of Pennsylvania natural gas, measured understandard conditions, were to be burned. How much water by weight wouldresult? CHAPTER XIX MOLECULAR WEIGHTS, ATOMIC WEIGHTS, FORMULAS ~Introduction. ~ In the chapter on The Atomic Theory, it was shown that ifit were true that two elements uniting to form a compound alwayscombined in the ratio of one atom of one element to one atom of theother element, it would be a very easy matter to decide upon figureswhich would represent the relative weights of the different atoms. Itwould only be necessary to select some one element as a standard anddetermine the weight of every element which combines with a definiteweight (say 1 g. ) of the standard element. The figures so obtained wouldevidently represent the relative weights of the atoms. But the law of multiple proportion at once reminds us that two elementsmay unite in several proportions; and there is no simple way todetermine the number of atoms present in the molecule of any compound. Consequently the problem of deciding upon the relative atomic weights isnot an easy one. To the solution of this problem we must now turn. ~Dalton's method of determining atomic weights. ~ When Dalton firstadvanced the atomic theory he attempted to solve this problem by verysimple methods. He thought that when only one compound of two elementsis known it is reasonable to suppose that it contains one atom of eachelement. He therefore gave the formula HO to water, and HN to ammonia. When more than two compounds were known he assumed that the mostfamiliar or the most stable one had the simple formula. He thendetermined the atomic weight as explained above. The results heobtained were contradictory and very far from satisfactory, and it wassoon seen that some other method, resting on much more scientificgrounds, must be found to decide what compounds, if any, have a singleatom of each element present. ~Determination of atomic weights. ~ Three distinct steps are involved inthe determination of the atomic weight of an element: (1) determinationof the equivalent, (2) determination of molecular weights of itscompounds, and (3) deduction of the exact atomic weight from theequivalent and molecular weights. ~1. Determination of the equivalent. ~ By the equivalent of an element ismeant the weight of the element which will combine with a fixed weightof some other element chosen as a standard. It has already beenexplained that oxygen has been selected as the standard element foratomic weights, with a weight of 16. This same standard will serve verywell as a standard for equivalents. _The equivalent of an element is theweight of the element which will combine with 16 g. Of oxygen. _ Thus 16g. Of oxygen combines with 16. 03 g. Of sulphur, 65. 4 g. Of zinc, 215. 86g. Of silver, 70. 9 g. Of chlorine. These figures, therefore, representthe equivalent weights of these elements. ~Relation of atomic weights to equivalents. ~ According to the atomictheory combination always takes place between whole numbers of atoms. Thus one atom unites with one other, or with two or three; or two atomsmay unite with three, or three with five, and so on. When oxygen combines with zinc the combination must be between definitenumbers of the two kinds of atoms. Experiment shows that these twoelements combine in the ratio of 16 g. Of oxygen to 65. 4 g. Of zinc. Ifone atom of oxygen combines with one atom of zinc, then this ratio mustbe the ratio between the weights of the two atoms. If one atom of oxygencombines with two atoms of zinc, then the ratio between the weights ofthe two atoms will be 16: 32. 7. If two atoms of oxygen combine with oneatom of zinc, the ratio by weight between the two atoms will be 8: 65. 4. It is evident, therefore, that the real atomic weight of an element mustbe some multiple or submultiple of the equivalent; in other words, theequivalent multiplied by 1/2, 1, 2, or 3 will give the atomic weight. ~Combining weights. ~ A very interesting relation holds good between theequivalents of the various elements. We have just seen that the figures16. 03, 65. 4, 215. 86, and 70. 9 are the equivalents respectively ofsulphur, zinc, silver, and chlorine. These same figures represent theratios by weight in which these elements combine among themselves. Thus215. 86 g. Of silver combine with 70. 9 g. Of chlorine and with 2 × 16. 03g. Of sulphur. 65. 4 g. Of zinc combine with 70. 9 g. Of chlorine and 2 ×16. 03 g. Of sulphur. By taking the equivalent or some multiple of it a value can be obtainedfor each element which will represent its combining value, and for thisreason is called its _combining weight_. It is important to notice thatthe fact that a combining weight can be obtained for each element is nota part of a theory, but is the direct result of experiment. ~Elements with more than one equivalent. ~ It will be remembered thatoxygen combines with hydrogen in two ratios. In one case 16 g. Of oxygencombine with 2. 016 g. Of hydrogen to form water; in the other 16 g. Ofoxygen combine with 1. 008 g. Of hydrogen to form hydrogen dioxide. Theequivalents of hydrogen are therefore 2. 016 and 1. 008. Barium combineswith oxygen in two proportions: in barium oxide the proportion is 16 g. Of oxygen to 137. 4 g. Of barium; in barium dioxide the proportion is 16g. Of oxygen to 68. 7 g. Of barium. In each case one equivalent is a simple multiple of the other, so thefact that there may be two equivalents does not add to the uncertainty. All we knew before was that the true atomic weight is some multiple ofthe equivalent. ~2. The determination of molecular weights. ~ To decide the question as towhich multiple of the equivalent correctly represents the atomic weightof an element, it has been found necessary to devise a method ofdetermining the molecular weights of compounds containing the element inquestion. Since the molecular weight of a compound is merely the sum ofthe weights of all the atoms present in it, it would seem to beimpossible to determine the molecular weight of a compound without firstknowing the atomic weights of the constituent atoms, and how many atomsof each element are present in the molecule. But certain facts have beendiscovered which suggest a way in which this can be done. ~Avogadro's hypothesis. ~ We have seen that the laws of Boyle, Charles, andGay-Lussac apply to all gases irrespective of their chemical character. This would lead to the inference that the structure of gases must bequite simple, and that it is much the same in all gases. In 1811 Avogadro, an Italian physicist, suggested that if we assume allgases under the same conditions of temperature and pressure to have thesame number of molecules in a given volume, we shall have a probableexplanation of the simplicity of the gas laws. It is difficult to provethe truth of this hypothesis by a simple experiment, but there are somany facts known which are in complete harmony with this suggestion thatthere is little doubt that it expresses the truth. Avogadro's hypothesismay be stated thus: _Equal volumes of all gases under the sameconditions of temperature and pressure contain the same number ofmolecules. _ ~Avogadro's hypothesis and molecular weights. ~ Assuming that Avogadro'shypothesis is correct, we have a very simple means for deciding upon therelative weights of molecules; for if equal volumes of two gases containthe same number of molecules, the weights of the two volumes must be inthe same ratio as the weights of the individual molecules which theycontain. If we adopt some one gas as a standard, we can express theweights of all other gases as compared with this one, and the samefigures will express the relative weights of the molecules of which thegases are composed. ~Oxygen as the standard. ~ It is important that the same standard should beadopted for the determination of molecular weights as has been decidedupon for atomic weights and equivalents, so that the three values may bein harmony with each other. Accordingly it is best to adopt oxygen asthe standard element with which to compare the molecular weights ofother gases, being careful to keep the oxygen atom equal to 16. ~The oxygen molecule contains two atoms. ~ One point must not beoverlooked, however. We desire to have our unit, the oxygen _atom_, equal to 16. The method of comparing the weights of gases just suggestedcompares the molecules of the gases with the _molecule_ of oxygen. Isthe molecule and the atom of oxygen the same thing? This question isanswered by the following considerations. We have seen that when steam is formed by the union of oxygen andhydrogen, two volumes of hydrogen combine with one volume of oxygen toform two volumes of steam. Let us suppose that the one volume of oxygencontains 100 molecules; then the two volumes of steam must, accordingto Avogadro's hypothesis, contain 200 molecules. But each of these 200molecules must contain at least one atom of oxygen, or 200 in all, andthese 200 atoms came from 100 molecules of oxygen. It follows that eachmolecule of oxygen must contain at least two atoms of oxygen. Evidently this reasoning merely shows that there are _at least_ twoatoms in the oxygen molecule. There may be more than that, but as thereis no evidence to this effect, we assume that the molecule contains twoatoms only. It is evident that if we wish to retain the value 16 for the atom ofoxygen we must take twice this value, or 32, for the value of the oxygenmolecule, when using it as a standard for molecular weights. ~Determination of the molecular weights of gases from their weightscompared with oxygen. ~ Assuming the molecular weight of oxygen to be 32, Avogadro's hypothesis gives us a ready means for determining themolecular weight of any other gas, for all that is required is to knowits weight compared with that of an equal volume of oxygen. For example, 1 l. Of chlorine is found by experiment to weigh 2. 216 times as much as1 l. Of oxygen. The molecular weight of chlorine must therefore be 2. 216× 32, or 70. 91. If, instead of comparing the relative weights of 1 l. Of the two gases, we select such a volume of oxygen as will weigh 32 g. , or the weight ingrams corresponding to the molecular weight of the gas, the calculationis much simplified. It has been found that 32 g. Of oxygen, understandard conditions, measure 22. 4 l. This same volume of hydrogen weighs2. 019 g. ; of chlorine 70. 9 g. ; of hydrochloric acid 36. 458 g. Theweights of these equal volumes must be proportional to their molecularweights, and since the weight of the oxygen is the same as the value ofits molecular weight, so too will the weights of the 22. 4 l. Of theother gases be equal to the value of their molecular weights. As a summary we can then make the following statement: _The molecularweight of any gas may be determined by calculating the weight of 22. 4 l. Of the gas, measured under standard conditions. _ ~Determination of molecular weights from density of gases. ~ In an actualexperiment it is easier to determine the density of a gas than theweight of a definite volume of it. The density of a gas is usuallydefined as its weight compared with that of an equal volume of air. Having determined the density of a gas, its weight compared with oxygenmay be determined by multiplying its density by the ratio between theweights of air and oxygen. This ratio is 0. 9046. To compare it with ourstandard for atomic weights we must further multiply it by 32, since thestandard is 1/32 the weight of oxygen molecules. The steps then arethese: 1. Determine the density of the gas (its weight compared with air). 2. Multiply by 0. 9046 to make the comparison with oxygen molecules. 3. Multiply by 32 to make the comparison with the unit for atomicweights. We have, then, the formula: molecular weight = density × 0. 9046 × 32; or, still more briefly, M. = D. × 28. 9. The value found by this method for the determination of molecularweights will of course agree with those found by calculating the weightof 22. 4 l. Of the gas, since both methods depend on the same principles. [Illustration: Fig. 69] ~Determination of densities of gases. ~ The relative weights of equal volumes of two gases can be easily determined. The following is one of the methods used. A small flask, such as is shown in Fig. 69, is filled with one of the gases, and after the temperature and pressure have been noted the flask is sealed up and weighed. The tip of the sealed end is then broken off, the flask filled with the second gas, and its weight determined. If the weight of the empty flask is subtracted from these two weighings, the relative weights of the gases is readily found. ~3. Deduction of atomic weights from molecular weights and equivalents. ~We have now seen how the equivalent of an element and the molecularweight of compounds containing the element can be obtained. Let us seehow it is possible to decide which multiple of the equivalent really isthe true atomic weight. As an example, let us suppose that theequivalent of nitrogen has been found to be 7. 02 and that it is desiredto obtain its atomic weight. The next step is to obtain the molecularweights of a large number of compounds containing nitrogen. Thefollowing will serve: ==================+============+=============+================+============== | | APPROXIMATE | PERCENTAGE OF | PART OF | DENSITY BY | MOLECULAR | NITROGEN BY | MOLECULAR | EXPERIMENT | WEIGHT | EXPERIMENT | WEIGHT DUE | | (D. × 28. 9) | | TO NITROGEN------------------+------------+-------------+----------------+--------------Nitrogen gas | 0. 9671 | 27. 95 | 100. 00 | 27. 95Nitrous oxide | 1. 527 | 44. 13 | 63. 70 | 27. 11Nitric oxide | 1. 0384 | 30. 00 | 46. 74 | 14. 02Nitrogen peroxide | 1. 580 | 45. 66 | 30. 49 | 13. 90Ammonia | 0. 591 | 17. 05 | 82. 28 | 14. 03Nitric acid | 2. 180 | 63. 06 | 22. 27 | 14. 03Hydrocyanic acid | 0. 930 | 26. 87 | 51. 90 | 13. 94==================+============+=============+================+============== ~Method of calculation. ~ The densities of the various gases in the firstcolumn of this table are determined by experiment, and are fairlyaccurate but not entirely so. By multiplying these densities by 28. 9 themolecular weights of the compounds as given in the second column areobtained. By chemical analysis it is possible to determine thepercentage composition of these substances, and the percentages ofnitrogen in them as determined by analysis are given in the thirdcolumn. If each of these molecular weights is multiplied in turn by thepercentage of nitrogen in the compound, the product will be the weightof the nitrogen in the molecular weight of the compound. This will bethe sum of the weights of the nitrogen atoms in the molecule. Thesevalues are given in the fourth column in the table. If a large number of compounds containing nitrogen are studied in thisway, it is probable that there will be included in the list at least onesubstance whose molecule contains a single nitrogen atom. In this casethe number in the fourth column will be the approximate atomic weight ofnitrogen. On comparing the values for nitrogen in the table it will beseen that a number which is approximately 14 is the smallest, and thatthe others are multiples of this. These compounds of higher value, therefore, contain more than one nitrogen atom in the molecule. ~Accurate determination of atomic weights. ~ Molecular weights cannot bedetermined very accurately, and consequently the part in them due tonitrogen is a little uncertain, as will be seen in the table. All we cantell by this method is that the true weight is very near 14. Theequivalent can however be determined very accurately, and we have seenthat it is some multiple or submultiple of the true atomic weight. Since molecular-weight determinations have shown that in the case ofnitrogen the atomic weight is near 14, and we have found the equivalentto be 7. 02, it is evident that the true atomic weight is twice theequivalent, or 7. 02 × 2 = 14. 04. ~Summary. ~ These, then, are the steps necessary to establish the atomicweight of an element. 1. Determine the equivalent accurately by analysis. 2. Determine the molecular weight of a large number of compounds of theelement, and by analysis the part of the molecular weight due to theelement. The smallest number so obtained will be approximately theatomic weight. 3. Multiply the equivalent by the small whole number (usually 1, 2, or3), which will make a number very close to the approximate atomicweight. The figure so obtained will be the true atomic weight. ~Molecular weights of the elements. ~ It will be noticed that the molecularweight of nitrogen obtained by multiplying its density by 28. 9 is 28. 08. Yet the atomic weight of nitrogen as deduced from a study of its gaseouscompounds is 14. 04. The simplest explanation that can be given for thisis that the gaseous nitrogen is made up of molecules, each of whichcontains two atoms. In this respect it resembles oxygen; for we haveseen that an entirely different line of reasoning leads us to believethat the molecule of oxygen contains two atoms. When we wish to indicatemolecules of these gases the symbols N_{2} and O_{2} should be used. When we desire to merely show the weights taking part in a reaction thisis not necessary. The vapor densities of many of the elements show that, like oxygen andnitrogen, their molecules consist of two atoms. In other cases, particularly among the metals, the molecule and the atom are identical. Still other elements have four atoms in their molecules. While oxygen contains two atoms in its molecules, a study of ozone hasled to the conclusion that it has three. The formation of ozone fromoxygen can therefore be represented by the equation 3O_{2} = 2O_{3}. ~Other methods of determining molecular weights. ~ It will be noticed thatAvogadro's law gives us a method by which we can determine the relativeweights of the molecules of two gases because it enables us to tell whenwe are dealing with an equal number of the two kinds of molecules. If byany other means we can get this information, we can make use of theknowledge so gained to determine the molecular weights of the twosubstances. ~Raoult's laws. ~ Two laws have been discovered which give us just suchinformation. They are known as Raoult's laws, and can be stated asfollows: 1. _When weights of substances which are proportional to their molecularweights are dissolved in the same weight of solvent, the rise of theboiling point is the same in each case. _ 2. _When weights of substances which are proportional to their molecularweights are dissolved in the same weight of solvent, the lowering of thefreezing point is the same in each case. _ By taking advantage of these laws it is possible to determine when twosolutions contain the same number of molecules of two dissolvedsubstances, and consequently the relative molecular weights of the twosubstances. ~Law of Dulong and Petit. ~ In 1819 Dulong and Petit discovered a veryinteresting relation between the atomic weight of an element and itsspecific heat, which holds true for elements in the solid state. Ifequal weights of two solids, say, lead and silver, are heated throughthe same range of temperature, as from 10° to 20°, it is found that verydifferent amounts of heat are required. The amount of heat required tochange the temperature of a solid or a liquid by a definite amountcompared with the amount required to change the temperature of an equalweight of water by the same amount is called its specific heat. Dulongand Petit discovered the following law: _The specific heat of an elementin the solid form multiplied by its atomic weight is approximately equalto the constant 6. 25. _ That is, at. Wt. × sp. Ht. = 6. 25. Consequently, 6. 25 at. Wt. = -------- sp. Ht. This law is not very accurate, but it is often possible by means of itto decide upon what multiple of the equivalent is the real atomicweight. Thus the specific heat of iron is found by experiment to be0. 112, and its equivalent is 27. 95. 6. 25 ÷ 0. 112 = 55. 8. We see, therefore, that the atomic weight is twice the equivalent, or 55. 9. ~How formulas are determined. ~ It will be well in connection withmolecular weights to consider how the formula of a compound is decidedupon, for the two subjects are very closely associated. Some exampleswill make clear the method followed. The molecular weight of a substance containing hydrogen and chlorine was36. 4. By analysis 36. 4 parts of the substance was found to contain 1part of hydrogen and 35. 4 parts of chlorine. As these are the simpleatomic weights of the two elements, the formula of the compound must beHCl. A substance consisting of oxygen and hydrogen was found to have amolecular weight of 34. Analysis showed that in 34 parts of thesubstance there were 2 parts of hydrogen and 32 parts of oxygen. Dividing these figures by the atomic weights of the two elements, we get2 ÷ 1 = 2 for H; 32 ÷ 16 = 2 for O. The formula is therefore H_{2}O_{2}. A substance containing 2. 04% H, 32. 6% S, and 65. 3% O was found to have amolecular weight of 98. In these 98 parts of the substance there are 98× 2. 04% = 2 parts of H, 98 × 32. 6% = 32 parts of S, and 98 × 65. 3% = 64parts of O. If the molecule weighs 98, the hydrogen atoms present musttogether weigh 2, the sulphur atoms 32, and the oxygen atoms 64. Dividing these figures by the respective atomic weights of the threeelements, we have, for H, 2 ÷ 1 = 2 atoms; for S, 32 ÷ 32 = 1 atom; forO, 64 ÷ 16 = 4 atoms. Hence the formula is H_{2}SO_{4}. We have, then, this general procedure: Find the percentage compositionof the substance and also its molecular weight. Multiply the molecularweight successively by the percentage of each element present, to findthe amount of the element in the molecular weight of the compound. Thefigures so obtained will be the respective parts of the molecular weightdue to the several atoms. Divide by the atomic weights of the respectiveelements, and the quotient will be the number of atoms present. ~Avogadro's hypothesis and chemical calculations. ~ This law simplifiesmany chemical calculations. 1. _Application to volume relations in gaseous reactions. _ Since equalvolumes of gases contain an equal number of molecules, it follows thatwhen an equal number of gaseous molecules of two or more gases take partin a reaction, the reaction will involve equal volumes of the gases. Inthe equation C_{2}H_{2}O_{4} = H_{2}O + CO_{2} + CO, since 1 molecule of each of the gases CO_{2} and CO is set free fromeach molecule of oxalic acid, the two substances must always be set freein equal volumes. Acetylene burns in accordance with the equation 2C_{2}H_{2} + 5O_{2} = 4CO_{2} + 2H_{2}O. Hence 2 volumes of acetylene will react with 5 volumes of oxygen to form4 volumes of carbon dioxide and 2 volumes of steam. That the volumerelations may be correct a gaseous element must be given its molecularformula. Thus oxygen must be written O_{2} and not 2O. 2. _Application to weights of gases. _ It will be recalled that themolecular weight of a gas is determined by ascertaining the weight of22. 4 l. Of the gas. This weight in grams is called the _gram-molecularweight_ of a gas. If the molecular weight of any gas is known, theweight of a liter of the gas under standard conditions may be determinedby dividing its gram-molecular weight by 22. 4. Thus the gram-molecularweight of a hydrochloric acid gas is 36. 458. A liter of the gas willtherefore weigh 36. 458 ÷ 22. 4 = 1. 627 g. EXERCISES 1. From the following data calculate the atomic weight of sulphur. Theequivalent, as obtained by an analysis of sulphur dioxide, is 16. 03. Thedensities and compositions of a number of compounds containing sulphurare as follows: NAME DENSITY COMPOSITION BY PERCENTAGEHydrosulphuric acid 1. 1791 S = 94. 11 H = 5. 89Sulphur dioxide 2. 222 S = 50. 05 O = 49. 95Sulphur trioxide 2. 74 S = 40. 05 O = 59. 95Sulphur chloride 4. 70 S = 47. 48 Cl = 52. 52Sulphuryl chloride 4. 64 S = 23. 75 Cl = 52. 53 O = 23. 70Carbon disulphide 2. 68 S = 84. 24 C = 15. 76 2. Calculate the formulas for compounds of the following compositions: MOLECULAR WEIGHT(1) S = 39. 07% O = 58. 49% H = 2. 44% 81. 0(2) Ca = 29. 40 S = 23. 56 O = 47. 04 136. 2(3) K = 38. 67 N = 13. 88 O = 47. 45 101. 2 3. The molecular weight of ammonia is 17. 06; of sulphur dioxide is64. 06; of chlorine is 70. 9. From the molecular weight calculate theweight of 1 l. Of each of these gases. Compare your results with thetable on the back cover of the book. 4. From the molecular weight of the same gases calculate the density ofeach, referred to air as a standard. 5. A mixture of 50 cc. Of carbon monoxide and 50 cc. Of oxygen wasexploded in a eudiometer, (a) What gases remained in the tube afterthe explosion? (b) What was the volume of each? 6. In what proportion must acetylene and oxygen be mixed to produce thegreatest explosion? 7. Solve Problem 18, Chapter XVII, without using molecular weights. Compare your results. 8. Solve Problem 10, Chapter XVIII, without using molecular weights. Compare your results. 9. The specific heat of aluminium is 0. 214; of lead is 0. 031. From thesespecific heats calculate the atomic weights of each of the elements. CHAPTER XX THE PHOSPHORUS FAMILY ================================================== | | ATOMIC | | MELTING | SYMBOL | WEIGHT | DENSITY | POINT-----------+--------+---------+---------+---------Phosphorus | P | 31. 0 | 1. 8 | 43. 3°Arsenic | As | 75. 0 | 5. 73 | ---Antimony | Sb | 120. 2 | 6. 7 | 432°Bismuth | Bi | 208. 5 | 9. 8 | 270°================================================== ~The family. ~ The elements constituting this family belong in the samegroup with nitrogen and therefore resemble it in a general way. Theyexhibit a regular gradation of physical properties, as is shown in theabove table. The same general gradation is also found in their chemicalproperties, phosphorus being an acid-forming element, while bismuth isessentially a metal. The other two elements are intermediate inproperties. ~Compounds. ~ In general the elements of the family form compounds havingsimilar composition, as is shown in the following table: PH_{3} PCl_{3} PCl_{5} P_{2}O_{3} P_{2}O_{5} AsH_{3} AsCl_{3} AsCl_{5} As_{2}O_{3} As_{2}O_{5} SbH_{3} SbCl_{3} SbCl_{5} Sb_{2}O_{3} Sb_{2}O_{5} . .. . BiCl_{3} BiCl_{5} Bi_{2}O_{3} Bi_{2}O_{5} In the case of phosphorus, arsenic, and antimony the oxides are acidanhydrides. Salts of at least four acids of each of these three elementsare known, the free acid in some instances being unstable. The relationof these acids to the corresponding anhydrides may be illustrated asfollows, phosphorus being taken as an example: P_{2}O_{3} + 3H_{2}O = 2H_{3}PO_{3} (phosphorous acid). P_{2}O_{5} + 3H_{2}O = 2H_{3}PO_{4} (phosphoric acid). P_{2}O_{5} + 2H_{2}O = H_{4}P_{2}O_{7} (pyrophosphoric acid). P_{2}O_{5} + H_{2}O = 2HPO_{3} (metaphosphoric acid). PHOSPHORUS ~History. ~ The element phosphorus was discovered by the alchemist Brand, of Hamburg, in 1669, while searching for the philosopher's stone. Owingto its peculiar properties and the secrecy which was maintained aboutits preparation, it remained a very rare and costly substance until thedemand for it in the manufacture of matches brought about its productionon a large scale. ~Occurrence. ~ Owing to its great chemical activity phosphorus never occursfree in nature. In the form of phosphates it is very abundant and widelydistributed. _Phosphorite_ and _sombrerite_ are mineral forms of calciumphosphate, while _apatite_ consists of calcium phosphate together withcalcium fluoride or chloride. These minerals form very large depositsand are extensively mined for use as fertilizers. Calcium phosphate is aconstituent of all fertile soil, having been supplied to the soil by thedisintegration of rocks containing it. It is the chief mineralconstituent of bones of animals, and bone ash is therefore nearly purecalcium phosphate. ~Preparation. ~ Phosphorus is now manufactured from bone ash or a puremineral phosphate by heating the phosphate with sand and carbon in anelectric furnace. The materials are fed in at M (Fig. 70) by the feedscrew F. The phosphorus vapor escapes at P and is condensed underwater, while the calcium silicate is tapped off as a liquid at S. Thephosphorus obtained in this way is quite impure, and is purified bydistillation. [Illustration: Fig. 70] ~Explanation of the reaction. ~ To understand the reaction which occurs, it must be remembered that a volatile acid anhydride is expelled from its salts when heated with an anhydride which is not volatile. Thus, when sodium carbonate and silicon dioxide are heated together the following reaction takes place: Na_{2}CO_{3} + SiO_{2} = Na_{2}SiO_{3} + CO_{2}. Silicon dioxide is a less volatile anhydride than phosphoric anhydride (P_{2}O_{5}), and when strongly heated with a phosphate the phosphoric anhydride is driven out, thus: Ca_{3}(PO_{4})_{2} + 3SiO_{2} = 3CaSiO_{3} + P_{2}O_{5}. If carbon is added before the heat is applied, the P_{2}O_{5} is reduced to phosphorus at the same time, according to the equation P_{2}O_{5} + 5C = 2P + 5CO. ~Physical properties. ~ The purified phosphorus is a pale yellowish, translucent, waxy solid which melts at 43. 3° and boils at 269°. It cantherefore be cast into any convenient form under warm water, and isusually sold in the market in the form of sticks. It is quite soft andcan be easily cut with a knife, but this must always be done while theelement is covered with water, since it is extremely inflammable, andthe friction of the knife blade is almost sure to set it on fire if cutin the air. It is not soluble in water, but is freely soluble in someother liquids, notably in carbon disulphide. Its density is 1. 8. ~Chemical properties. ~ Exposed to the air phosphorus slowly combines withoxygen, and in so doing emits a pale light, or phosphorescence, whichcan be seen only in a dark place. The heat of the room may easily raisethe temperature to the kindling point of phosphorus, when it burns witha sputtering flame, giving off dense fumes of oxide of phosphorus. Itburns with dazzling brilliancy in oxygen, and combines directly withmany other elements, especially with sulphur and the halogens. Onaccount of its great affinity for oxygen it is always preserved underwater. Phosphorus is very poisonous, from 0. 2 to 0. 3 gram being a fatal dose. Ground up with flour and water or similar substances, it is often usedas a poison for rats and other vermin. ~Precaution. ~ The heat of the body is sufficient to raise phosphorus above its kindling temperature, and for this reason it should always be handled with forceps and never with the bare fingers. Burns occasioned by it are very painful and slow in healing. ~Red phosphorus. ~ On standing, yellow phosphorus gradually undergoes aremarkable change, being converted into a dark red powder which has adensity of 2. 1. It no longer takes fire easily, neither does it dissolvein carbon disulphide. It is not poisonous and, in fact, seems to be anentirely different substance. The velocity of this change increases withrise in temperature, and the red phosphorus is therefore prepared byheating the yellow just below the boiling point (250°-300°). Whendistilled and quickly condensed the red form changes back to the yellow. This is in accordance with the general rule that when a substancecapable of existing in several allotropic forms is condensed from a gasor crystallized from the liquid state, the more unstable variety formsfirst, and this then passes into the more stable forms. ~Matches. ~ The chief use of phosphorus is in the manufacture of matches. Common matches are made by first dipping the match sticks into some inflammable substance, such as melted paraffin, and afterward into a paste consisting of (1) phosphorus, (2) some oxidizing substance, such as manganese dioxide or potassium chlorate, and (3) a binding material, usually some kind of glue. On friction the phosphorus is ignited, the combustion being sustained by the oxidizing agent and communicated to the wood by the burning paraffin. In sulphur matches the paraffin is replaced by sulphur. In safety matches _red_ phosphorus, an oxidizing agent, and some gritty material such as emery is placed on the side of the box, while the match tip is provided as before with an oxidizing agent and an easily oxidized substance, usually antimony sulphide. The match cannot be ignited easily by friction, save on the prepared surface. ~Compounds of phosphorus with hydrogen. ~ Phosphorus forms severalcompounds with hydrogen, the best known of which is phosphine (PH_{3})analogous to ammonia (NH_{3}). ~Preparation of phosphine. ~ Phosphine is usually made by heatingphosphorus with a strong solution of potassium hydroxide, the reactionbeing a complicated one. [Illustration: Fig. 71] The experiment can be conveniently made in the apparatus shown in Fig. 71. A strong solution of potassium hydroxide together with several small bits of phosphorus are placed in the flask A, and a current of coal gas is passed into the flask through the tube B until all the air has been displaced. The gas is then turned off and the flask is heated. Phosphine is formed in small quantities and escapes through the delivery tube, the exit of which is just covered by the water in the vessel C. Each bubble of the gas as it escapes into the air takes fire, and the product of combustion (P_{2}O_{5}) forms beautiful small rings, which float unbroken for a considerable time in quiet air. The pure phosphine does not take fire spontaneously. When prepared as directed above, impurities are present which impart this property. ~Properties. ~ Phosphine is a gas of unpleasant odor and is exceedinglypoisonous. Like ammonia it forms salts with the halogen acids. Thus wehave phosphonium chloride (PH_{4}Cl) analogous to ammonium chloride(NH_{4}Cl). The phosphonium salts are of but little importance. ~Oxides of phosphorus. ~ Phosphorus forms two well-known oxides, --thetrioxide (P_{2}O_{3}) and the pentoxide (P_{2}O_{5}), sometimes calledphosphoric anhydride. When phosphorus burns in an insufficient supply ofair the product is partially the trioxide; in oxygen or an excess of airthe pentoxide is formed. The pentoxide is much the better known of thetwo. It is a snow-white, voluminous powder whose most marked property isits great attraction for water. It has no chemical action upon mostgases, so that they can be very thoroughly dried by allowing them topass through properly arranged vessels containing phosphorus pentoxide. ~Acids of phosphorus. ~ The important acids of phosphorus are thefollowing: H_{3}PO_{3} phosphorous acid. H_{3}PO_{4} phosphoric acid. H_{4}P_{2}O_{7} pyrophosphoric acid. HPO_{3} metaphosphoric acid. These may be regarded as combinations of the oxides of phosphorus withwater according to the equations given in the discussion of thecharacteristics of the family. 1. _Phosphorous acid_ (H_{3}PO_{3}). Neither the acid nor its salts areat all frequently met with in chemical operations. It can be easilyobtained, however, in the form of transparent crystals when phosphorustrichloride is treated with water and the resulting solution isevaporated: PCl_{3} + 3H_{2}O = H_{3}PO_{3} + 3HCl. Its most interesting property is its tendency to take up oxygen and passover into phosphoric acid. 2. _Orthophosphoric acid (phosphoric acid)_ (H_{3}PO_{4}). This acid canbe obtained by dissolving phosphorus pentoxide in boiling water, asrepresented in the equation P_{2}O_{5} + 3H_{2}O = 2H_{3}PO_{4}. It is usually made by treating calcium phosphate with concentratedsulphuric acid. The calcium sulphate produced in the reaction is nearlyinsoluble, and can be filtered off, leaving the phosphoric acid insolution. Very pure acid is made by oxidizing phosphorus with nitricacid. It forms large colorless crystals which are exceedingly soluble inwater. Being a tribasic acid, it forms acid as well as normal salts. Thus the following compounds of sodium are known: NaH_{2}PO_{4} monosodium hydrogen phosphate. Na_{2}HPO_{4} disodium hydrogen phosphate. Na_{3}PO_{4} normal sodium phosphate. These salts are sometimes called respectively primary, secondary, andtertiary phosphates. They may be prepared by bringing togetherphosphoric acid and appropriate quantities of sodium hydroxide. Phosphoric acid also forms mixed salts, that is, salts containing twodifferent metals. The most familiar compound of this kind is microcosmicsalt, which has the formula Na(NH_{4})HPO_{4}. _Orthophosphates. _ The orthophosphates form an important class of salts. The normal salts are nearly all insoluble and many of them occur innature. The secondary phosphates are as a rule insoluble, while most ofthe primary salts are soluble. 3. _Pyrophosphoric acid_ (H_{4}P_{2}O_{7}). On heating orthophosphoricacid to about 225° pyrophosphoric acid is formed in accordance with thefollowing equation: 2H_{3}PO_{4} = H_{4}P_{2}O_{7} + H_{2}O. It is a white crystalline solid. Its salts can be prepared by heating asecondary phosphate: 2Na_{2}HPO_{4} = Na_{4}P_{2}O_{7} + H_{2}O. 4. _Metaphosphoric acid (glacial phosphoric acid)_ (HPO_{3}). This acidis formed when orthophosphoric acid is heated above 400°: H_{3}PO_{4} = HPO_{3} + H_{2}O. It is also formed when phosphorus pentoxide is treated with cold water: P_{2}O_{5} + H_{2}O = 2HPO_{3}. It is a white crystalline solid, and is so stable towards heat that itcan be fused and even volatilized without decomposition. On cooling fromthe fused state it forms a glassy solid, and on this account is oftencalled glacial phosphoric acid. It possesses the property of dissolvingsmall quantities of metallic oxides, with the formation of compoundswhich, in the case of certain metals, have characteristic colors. It istherefore used in the detection of these metals. While the secondary phosphates, on heating, give salts of pyrophosphoricacid, the primary phosphates yield salts of metaphosphoric acid. Theequations representing these reactions are as follows: 2Na_{2}HPO_{4} = Na_{4}P_{3}O_{7} + H_{2}O, NaH_{2}PO_{4} = NaPO_{3} + H_{2}O. ~Fertilizers. ~ When crops are produced year after year on the same fieldcertain constituents of the soil essential to plant growth are removed, and the soil becomes impoverished and unproductive. To make the landonce more fertile these constituents must be replaced. The calciumphosphate of the mineral deposits or of bone ash serves well as amaterial for restoring phosphorus to soils exhausted of that essentialelement; but a more soluble substance, which the plants can more readilyassimilate, is desirable. It is better, therefore, to convert theinsoluble calcium phosphate into the soluble primary phosphate before itis applied as fertilizer. It will be seen by reference to the formulasfor the orthophosphates (see page 244) that in a primary phosphate onlyone hydrogen atom of phosphoric acid is replaced by a metal. Since thecalcium atom always replaces two hydrogen atoms, it might be thoughtthat there could be no primary calcium phosphate; but if the calciumatom replaces one hydrogen atom from each of two molecules of phosphoricacid, the salt Ca(H_{2}PO_{4})_{2} will result, and this is a primaryphosphate. It can be made by treatment of the normal phosphate with thenecessary amount of sulphuric acid, calcium sulphate being formed at thesame time, thus: Ca_{3}(PO_{4})_{2} + 2H_{2}SO_{4} = Ca(H_{2}PO_{4})_{2} + 2CaSO_{4}. The resulting mixture is a powder, which is sold as a fertilizer underthe name of "superphosphate of lime. " ARSENIC ~Occurrence. ~ Arsenic occurs in considerable quantities in nature as thenative element, as the sulphides realgar (As_{2}S_{2}) and orpiment(As_{2}S_{3}), as oxide (As_{2}O_{3}), and as a constituent of manymetallic sulphides, such as arsenopyrite (FeAsS). ~Preparation. ~ The element is prepared by purifying the native arsenic, orby heating the arsenopyrite in iron tubes, out of contact with air, when the reaction expressed by the following equation occurs: FeAsS = FeS + As. The arsenic, being volatile, condenses in chambers connected with theheated tubes. It is also made from the oxide by reduction with carbon: 2As_{2}O_{3} + 3C = 4As + 3CO_{2}. ~Properties. ~ Arsenic is a steel-gray, metallic-looking substance ofdensity 5. 73. Though resembling metals in appearance, it is quitebrittle, being easily powdered in a mortar. When strongly heated itsublimes, that is, it passes into a vapor without melting, and condensesagain to a crystalline solid when the vapor is cooled. Like phosphorusit can be obtained in several allotropic forms. It alloys readily withsome of the metals, and finds its chief use as an alloy with lead, whichis used for making shot, the alloy being harder than pure lead. Whenheated on charcoal with the blowpipe it is converted into an oxide whichvolatilizes, leaving the charcoal unstained by any oxide coating. Itburns readily in chlorine gas, forming arsenic trichloride, -- As + 3Cl = AsCl_{3}. Unlike most of its compounds, the element itself is not poisonous. ~Arsine~ (AsH_{3}). When any compound containing arsenic is brought intothe presence of nascent hydrogen, arsine (AsH_{3}), corresponding tophosphine and ammonia, is formed. The reaction when oxide of arsenic isso treated is As_{2}O_{3} + 12H = 2AsH_{3} + 3H_{2}O. Arsine is a gas with a peculiar garlic-like odor, and is intenselypoisonous. A single bubble of pure gas has been known to prove fatal. Itis an unstable compound, decomposing into its elements when heated to amoderate temperature. It is combustible, burning with a palebluish-white flame to form arsenic trioxide and water when air is inexcess: 2AsH_{3} + 6O = As_{2}O_{3} + 3H_{2}O. When the supply of air is deficient water and metallic arsenic areformed: 2AsH_{3} + 3O = 3H_{2}O + 2As. These reactions make the detection of even minute quantities of arsenica very easy problem. [Illustration: Fig. 72] ~Marsh's test for arsenic. ~ The method devised by Marsh for detecting arsenic is most frequently used, the apparatus being shown in Fig. 72. Hydrogen is generated in the flask A by the action of dilute sulphuric acid on zinc, is dried by passing over calcium chloride in the tube B, and after passing through the hard-glass tube C is ignited at the jet D. If a substance containing arsenic is now introduced into the generator A, the arsenic is converted into arsine by the action of the nascent hydrogen, and passes to the jet along with the hydrogen. If the tube C is strongly heated at some point near the middle, the arsine is decomposed while passing this point and the arsenic is deposited just beyond the heated point in the form of a shining, brownish-black mirror. If the tube is not heated, the arsine burns along with the hydrogen at the jet. Under these conditions a small porcelain dish crowded down into the flame is blackened by a spot of metallic arsenic, for the arsine is decomposed by the heat of the flame, and the arsenic, cooled below its kindling temperature by the cold porcelain, deposits upon it as a black spot. Antimony conducts itself in the same way as arsenic, but the antimony deposit is more sooty in appearance. The two can also be distinguished by the fact that sodium hypochlorite (NaClO) dissolves the arsenic deposit, but not that formed by antimony. ~Oxides of arsenic. ~ Arsenic forms two oxides, As_{2}O_{3} andAs_{2}O_{5}, corresponding to those of phosphorus. Of these arseniousoxide, or arsenic trioxide (As_{2}O_{3}), is much better known, and isthe substance usually called white arsenic, or merely arsenic. It isfound as a mineral, but is usually obtained as a by-product in burningpyrite in the sulphuric-acid industry. The pyrite has a small amount ofarsenopyrite in it, and when this is burned arsenious oxide is formed asa vapor together with sulphur dioxide: 2FeAsS + 10O = Fe_{2}O_{3} + As_{2}O_{3} + 2SO_{2}. The arsenious oxide is condensed in appropriate chambers. It is a ratherheavy substance, obtained either as a crystalline powder or as large, vitreous lumps, resembling lumps of porcelain in appearance. It is verypoisonous, from 0. 2 to 0. 3 g. Being a fatal dose. It is frequently givenas a poison, since it is nearly tasteless and does not act very rapidly. This slow action is due to the fact that it is not very soluble, andhence is absorbed slowly by the system. Arsenious oxide is also used asa chemical reagent in glass making and in the dye industry. ~Acids of arsenic. ~ Like the corresponding oxides of phosphorus, theoxides of arsenic are acid anhydrides. In solution they combine withbases to form salts, corresponding to the salts of the acids ofphosphorus. Thus we have salts of the following acids: H_{3}AsO_{3} arsenious acid. H_{3}AsO_{4} orthoarsenic acid. H_{4}As_{2}O_{3} pyroarsenic acid. HAsO_{3} metarsenic acid. Several other acids of arsenic are also known. Not all of these can beobtained as free acids, since they tend to lose water and form theoxides. Thus, instead of obtaining arsenious acid (H_{3}AsO_{3}), theoxide As_{2}O_{3} is obtained: 2H_{3}AsO_{3} = As_{2}O_{3} + 3H_{2}O. Salts of all the acids are known, however, and some of them havecommercial value. Most of them are insoluble, and some of the coppersalts, which are green, are used as pigments. Paris green, which has acomplicated formula, is a well-known insecticide. ~Antidote for arsenical poisoning. ~ The most efficient antidote forarsenic poisoning is ferric hydroxide. It is prepared as needed, according to the equation Fe_{2}(SO_{4})_{3} + 3Mg(OH)_{2} = 2Fe(OH)_{3} + 3MgSO_{4}. ~Sulphides of arsenic. ~ When hydrogen sulphide is passed into an acidifiedsolution containing an arsenic compound the arsenic is precipitated as abright yellow sulphide, thus: 2H_{3}AsO_{3} + 3H_{2}S = As_{2}S_{3} + 6H_{2}O, 2H_{3}AsO_{4} + 5H_{2}S = As_{2}S_{5} + 8H_{2}O. In this respect arsenic resembles the metallic elements, many of whichproduce sulphides under similar conditions. The sulphides of arsenic, both those produced artificially and those found in nature, are used asyellow pigments. ANTIMONY ~Occurrence. ~ Antimony occurs in nature chiefly as the sulphide(Sb_{2}S_{3}), called stibnite, though it is also found as oxide and asa constituent of many complex minerals. ~Preparation. ~ Antimony is prepared from the sulphide in a very simplemanner. The sulphide is melted with scrap iron in a furnace, when theiron combines with the sulphur to form a slag, or liquid layer of meltediron sulphide, while the heavier liquid, antimony, settles to the bottomand is drawn off from time to time. The reaction involved is representedby the equation Sb_{2}S_{3} + 3Fe = 2Sb + 3FeS. ~Physical properties. ~ Antimony is a bluish-white, metallic-lookingsubstance whose density is 6. 7. It is highly crystalline, hard, and verybrittle. It has a rather low melting point (432°) and expands verynoticeably on solidifying. ~Chemical properties. ~ In chemical properties antimony resembles arsenicin many particulars. It forms the oxides Sb_{2}O_{3} and Sb_{2}O_{5}, and in addition Sb_{2}O_{4}. It combines with the halogen elements withgreat energy, burning brilliantly in chlorine to form antimonytrichloride (SbCl_{3}). When heated on charcoal with the blowpipe it isoxidized and forms a coating of antimony oxide on the charcoal which hasa characteristic bluish-white color. ~Stibine~ (SbH_{3}). The gas stibine (SbH_{3}) is formed under conditionswhich are very similar to those which produce arsine, and it closelyresembles the latter compound, though it is still less stable. It isvery poisonous. ~Acids of antimony. ~ The oxides Sb_{2}O_{3} and Sb_{2}O_{5} are weak acid anhydrides and are capable of forming two series of acids corresponding in formulas to the acids of phosphorus and arsenic. They are much weaker, however, and are of little practical importance. ~Sulphides of antimony. ~ Antimony resembles arsenic in that hydrogen sulphide precipitates it as a sulphide when conducted into an acidified solution containing an antimony compound: 2SbCl_{3} + 3H_{2}S = Sb_{2}S_{3} + 6HCl, 2SbCl_{5} + 5H_{2}S = Sb_{2}S_{5} + 10HCl. The two sulphides of antimony are called the trisulphide and the pentasulphide respectively. When prepared in this way they are orange-colored substances, though the mineral stibnite is black. ~Metallic properties of antimony. ~ The physical properties of the elementare those of a metal, and the fact that its sulphide is precipitated byhydrogen sulphide shows that it acts like a metal in a chemical way. Many other reactions show that antimony has more of the properties of ametal than of a non-metal. The compound Sb(OH)_{3}, corresponding toarsenious acid, while able to act as a weak acid is also able to act asa weak base with strong acids. For example, when treated withconcentrated hydrochloric acid antimony chloride is formed: Sb(OH)_{3} + 3HCl = SbCl_{3} + 3H_{2}O. A number of elements act in this same way, their hydroxides under someconditions being weak acids and under others weak bases. ALLOYS Some metals when melted together thoroughly intermix, and on coolingform a homogeneous, metallic-appearing substance called an _alloy_. Notall metals will mix in this way, and in some cases definite chemicalcompounds are formed and separate out as the mixture solidifies, thusdestroying the uniform quality of the alloy. In general the meltingpoint of the alloy is below the average of the melting points of itsconstituents, and it is often lower than any one of them. Antimony forms alloys with many of the metals, and its chief commercialuse is for such purposes. It imparts to its alloys high density, ratherlow melting point, and the property of expanding on solidification. Such an alloy is especially useful in type founding, where fine linesare to be reproduced on a cast. Type metal consists of antimony, lead, and tin. Babbitt metal, used for journal bearings in machinery, containsthe same metals in a different proportion together with a smallpercentage of copper. BISMUTH ~Occurrence. ~ Bismuth is usually found in the uncombined form in nature. It also occurs as oxide and sulphide. Most of the bismuth of commercecomes from Saxony, and from Mexico and Colorado, but it is not anabundant element. ~Preparation. ~ It is prepared by merely heating the ore containing thenative bismuth and allowing the melted metal to run out into suitablevessels. Other ores are converted into oxides and reduced by heatingwith carbon. ~Physical properties. ~ Bismuth is a heavy, crystalline, brittle metalnearly the color of silver, but with a slightly rosy tint whichdistinguishes it from other metals. It melts at a low temperature (270°)and has a density of 9. 8. It is not acted upon by the air at ordinarytemperatures. ~Chemical properties. ~ When heated with the blowpipe on charcoal, bismuthgives a coating of the oxide Bi_{2}O_{3}. This has a yellowish-browncolor which easily distinguishes it from the oxides formed by othermetals. It combines very readily with the halogen elements, powderedbismuth burning readily in chlorine. It is not very easily acted upon byhydrochloric acid, but nitric and sulphuric acids act upon it in thesame way that they do upon copper. ~Uses. ~ Bismuth finds its chief use as a constituent of alloys, particularly in those of low melting point. Some of these melt in hotwater. For example, Wood's metal, consisting of bismuth, lead, tin, andcadmium, melts at 60. 5°. ~Compounds of bismuth. ~ Unlike the other elements of this group, bismuthhas almost no acid properties. Its chief oxide, Bi_{2}O_{3}, is basic inits properties. It dissolves in strong acids and forms salts of bismuth: Bi_{2}O_{3} + 6HCl = 2BiCl_{3} + 3H_{2}O, Bi_{2}O_{3} + 6HNO_{3} = 2Bi(NO_{3})_{3} + 3H_{2}O. The nitrate and chloride of bismuth can be obtained as well-formedcolorless crystals. When treated with water the salts are decomposed inthe manner explained in the following paragraph. HYDROLYSIS Many salts such as those of antimony and bismuth form solutions whichare somewhat acid in reaction, and must therefore contain hydrogen ions. This is accounted for by the same principle suggested to explain thefact that solutions of potassium cyanide are alkaline in reaction (p. 210). Water forms an appreciable number of hydrogen and hydroxyl ions, and very weak bases such as bismuth hydroxide are dissociated to but avery slight extent. When Bi^{+++} ions from bismuth chloride, whichdissociates very readily, are brought in contact with the OH^{-} ionsfrom water, the two come to the equilibrium expressed in the equation Bi^{+++} + 3OH^{-} Bi(OH)_{3}. For every hydroxyl ion removed from the solution in this way a hydrogenion is left free, and the solution becomes acid in reaction. Reactions of this kind and that described under potassium cyanide arecalled _hydrolysis_. DEFINITION: _Hydrolysis is the action of water upon a salt to form anacid and a base, one of which is very slightly dissociated. _ ~Conditions favoring hydrolysis. ~ While hydrolysis is primarily due to theslight extent to which either the acid or the base formed isdissociated, several other factors have an influence upon the extent towhich it will take place. 1. _Influence of mass. _ Since hydrolysis is a reversible reaction, therelative masses of the reacting substances influence the point at whichequilibrium will be reached. In the equilibrium BiCl_{3} + 3H_{2}O Bi(OH)_{3} + 3HCl the addition of more water will result in the formation of more bismuthhydroxide and hydrochloric acid. The addition of more hydrochloric acidwill convert some of the bismuth hydroxide into bismuth chloride. 2. _Formation of insoluble substances. _ When one of the products ofhydrolysis is nearly insoluble in water the solution will becomesaturated with it as soon as a very little has been formed. All inexcess of this will precipitate, and the reaction will go on until theacid set free increases sufficiently to bring about an equilibrium. Thusa considerable amount of bismuth and antimony hydroxides areprecipitated when water is added to the chlorides of these elements. Thegreater the dilution the more hydroxide precipitates. The addition ofhydrochloric acid in considerable quantity will, however, redissolve theprecipitate. ~Partial hydrolysis. ~ In many cases the hydrolysis of a salt is onlypartial, resulting in the formation of basic salts instead of the freebase. Most of these basic salts are insoluble in water, which accountsfor their ready formation. Thus bismuth chloride may hydrolyze bysuccessive steps, as shown in the equations BiCl_{3} + H_{2}O = Bi(OH)Cl_{2} + HCl, BiCl_{3} + 2H_{2}O = Bi(OH)_{2}Cl + 2HCl, BiCl_{3} + 3H_{2}O = Bi(OH)_{3} + 3HCl. The basic salt so formed may also lose water, as shown in the equation Bi(OH)_{2}Cl = BiOCl + H_{2}O. The salt represented in the last equation is sometimes called bismuthoxychloride, or bismuthyl chloride. The corresponding nitrate, BiONO_{3}, is largely used in medicine under the name of subnitrate ofbismuth. In these two compounds the group of atoms, BiO, acts as aunivalent metallic radical and is called _bismuthyl_. Similar basicsalts are formed by the hydrolysis of antimony salts. EXERCISES 1. Name all the elements so far studied which possess allotropic forms. 2. What compounds would you expect phosphorus to form with bromine andiodine? Write the equations showing the action of water on thesecompounds. 3. In the preparation of phosphine, why is coal gas passed into theflask? What other gases would serve the same purpose? 4. Give the formula for the salt which phosphine forms with hydriodicacid. Give the name of the compound. 5. Could phosphoric acid be substituted for sulphuric acid in thepreparation of the common acids? 6. Write the equations for the preparation of the three sodium salts oforthophosphoric acid. 7. Why does a solution of disodium hydrogen phosphate react alkaline? 8. On the supposition that bone ash is pure calcium phosphate, whatweight of it would be required in the preparation of 1 kg. Ofphosphorus? 9. If arsenopyrite is heated in a current of air, what products areformed? 10. (a) Write equations for the complete combustion of hydrosulphuricacid, methane, and arsine. (b) In what respects are the reactionssimilar? 11. Write the equations for all the reactions involved in Marsh's testfor arsenic. 12. Write the names and formulas for the acids of antimony. 13. Write the equations showing the hydrolysis of antimony trichloride;of bismuth nitrate. 14. In what respects does nitrogen resemble the members of thephosphorus family? CHAPTER XXI SILICON, TITANIUM, BORON ================================================================= | | | | | | SYMBOL | ATOMIC | DENSITY | CHLORIDES | OXIDES | | WEIGHT | | |____________|________|________|_________|___________|____________ | | | | |Silicon | Si | 28. 4 | 2. 35 | SiCl_{4} | SiO_{2}Titanium | Ti | 48. 1 | 3. 5 | TiCl_{4} | TiO_{2}Boron | B | 11. 0 | 2. 45 | BCl_{3} | B_{2}O_{3}================================================================= ~General. ~ Each of the three elements, silicon, titanium, and boron, belongs to a separate periodic family, but they occur near together inthe periodic grouping and are very similar in both physical and chemicalproperties. Since the other elements in their families are either sorare that they cannot be studied in detail, or are best understood inconnection with other elements, it is convenient to consider these threetogether at this point. The three elements are very difficult to obtain in the free state, owingto their strong attraction for other elements. They can be prepared bythe action of aluminium or magnesium on their oxides and in impure stateby reduction with carbon in an electric furnace. They are very hard andmelt only at the highest temperatures. At ordinary temperatures they arenot attacked by oxygen, but when strongly heated they burn with greatbrilliancy. Silicon and boron are not attacked by acids under ordinaryconditions; titanium is easily dissolved by them. SILICON ~Occurrence. ~ Next to oxygen silicon is the most abundant element. It doesnot occur free in nature, but its compounds are very abundant and of thegreatest importance. It occurs almost entirely in combination withoxygen as silicon dioxide (SiO_{2}), often called silica, or with oxygenand various metals in the form of salts of silicic acids, or silicates. These compounds form a large fraction of the earth's crust. Most plantsabsorb small amounts of silica from the soil, and it is also found inminute quantities in animal organisms. ~Preparation. ~ The element is most easily prepared by reducing purepowdered quartz with magnesium powder: SiO_{2} + 2Mg = 2MgO + Si. ~Properties. ~ As would be expected from its place in the periodic table, silicon resembles carbon in many respects. It can be obtained in severalallotropic forms, corresponding to those of carbon. The crystallizedform is very hard, and is inactive toward reagents. The amorphousvariety has, in general, properties more similar to charcoal. ~Compounds of silicon with hydrogen and the halogens. ~ Silicon hydride(SiH_{4}) corresponds in formula to methane (CH_{4}), but its propertiesare more like those of phosphine (PH_{3}). It is a very inflammable gasof disagreeable odor, and, as ordinarily prepared, takes firespontaneously on account of the presence of impurities. Silicon combines with the elements of the chlorine family to form suchcompounds as SiCl_{4} and SiF_{4}. Of these silicon fluoride is the mostfamiliar and interesting. As stated in the discussion of fluorine, it isformed when hydrofluoric acid acts upon silicon dioxide or a silicate. With silica the reaction is thus expressed: SiO_{2} + 4HF = SiF_{4} + 2H_{2}O. It is a very volatile, invisible, poisonous gas. In contact with waterit is partially decomposed, as shown in the equation SiF_{4} + 4H_{2}O = 4HF + Si(OH)_{4}. The hydrofluoric acid so formed combines with an additional amount ofsilicon fluoride, forming the complex fluosilicic acid (H_{2}SiF_{6}), thus: 2HF + SiF_{4} = H_{2}SiF_{6}. ~Silicides. ~ As the name indicates, silicides are binary compoundsconsisting of silicon and some other element. They are very stable athigh temperatures, and are usually made by heating the appropriatesubstances in an electric furnace. The most important one is_carborundum_, which is a silicide of carbon of the formula CSi. It ismade by heating coke and sand, which is a form of silicon dioxide, in anelectric furnace, the process being extensively carried on at NiagaraFalls. The following equation represents the reaction SiO_{2} + 3C = CSi + 2CO. The substance so prepared consists of beautiful purplish-black crystals, which are very hard. Carborundum is used as an abrasive, that is, as amaterial for grinding and polishing very hard substances. Ferrosiliconis a silicide of iron alloyed with an excess of iron, which findsextensive use in the manufacture of certain kinds of steel. ~Manufacture of carborundum. ~ The mixture of materials is heated in alarge resistance furnace for about thirty-six hours. After the reactionis completed there is left a core of graphite G. Surrounding this coreis a layer of crystallized carborundum C, about 16 in. Thick. Outsidethis is a shell of amorphous carborundum A. The remaining materialsM are unchanged and are used for a new charge. [Illustration: Fig. 73] ~Silicon dioxide~ (_silica_) (SiO_{2}). This substance is found in a greatvariety of forms in nature, both in the amorphous and in the crystallinecondition. In the form of quartz it is found in beautifully formedsix-sided prisms, sometimes of great size. When pure it is perfectlytransparent and colorless. Some colored varieties are given specialnames, as amethyst (violet), rose quartz (pale pink), smoky or milkyquartz (colored and opaque). Other varieties of silicon dioxide, some ofwhich also contain water, are chalcedony, onyx, jasper, opal, agate, andflint. Sand and sandstone are largely silicon dioxide. ~Properties. ~ As obtained by chemical processes silicon dioxide is anamorphous white powder. In the crystallized state it is very hard andhas a density of 2. 6. It is insoluble in water and in most chemicalreagents, and requires the hottest oxyhydrogen flame for fusion. Acids, excepting hydrofluoric acid, have little action on it, and it requiresthe most energetic reducing agents to deprive it of oxygen. It is theanhydride of an acid, and consequently it dissolves in fused alkalis toform silicates. Being nonvolatile, it will drive out most otheranhydrides when heated to a high temperature with their salts, especially when the silicates so formed are fusible. The followingequations illustrate this property: Na_{2}CO_{3} + SiO_{2} = Na_{2}SiO_{3} + CO_{2}, Na_{2}SO_{4} + SiO_{2} = Na_{2}SiO_{3} + SO_{3}. ~Silicic acids. ~ Silicon forms two simple acids, orthosilicic acid(H_{4}SiO_{4}) and metasilicic acid (H_{2}SiO_{3}). Orthosilicic acid isformed as a jelly-like mass when orthosilicates are treated with strongacids such as hydrochloric. On attempting to dry this acid it loseswater, passing into metasilicic or common silicic acid: H_{4}SiO_{4} = H_{2}SiO_{3} + H_{2}O. Metasilicic acid when heated breaks up into silica and water, thus: H_{2}SiO_{3} = H_{2}O + SiO_{2}. ~Salts of silicic acids, --silicates. ~ A number of salts of the orthosilicic and metasilicic acids occur in nature. Thus mica (KAlSiO_{4}) is a salt of orthosilicic acid. ~Polysilicic acids. ~ Silicon has the power to form a great many complexacids which may be regarded as derived from the union of severalmolecules of the orthosilicic acid, with the loss of water. Thus we have 3H_{4}SiO_{4} = H_{4}Si_{3}O_{8} + 4H_{2}O. These acids cannot be prepared in the pure state, but their salts formmany of the crystalline rocks in nature. Feldspar, for example, has theformula KAlSi_{3}O_{8}, and is a mixed salt of the acidH_{4}Si_{3}O_{8}, whose formation is represented in the equation above. Kaolin has the formula Al_{2}Si_{2}O_{7}·2H_{2}O. Many other exampleswill be met in the study of the metals. ~Glass. ~ When sodium and calcium silicates, together with silicon dioxide, are heated to a very high temperature, the mixture slowly fuses to atransparent liquid, which on cooling passes into the solid called glass. Instead of starting with sodium and calcium silicates it is moreconvenient and economical to heat sodium carbonate (or sulphate) andlime with an excess of clean sand, the silicates being formed during theheating: Na_{2}CO_{3} + SiO_{2} = Na_{2}SiO_{3} + CO_{2}, CaO + SiO_{2} = CaSiO_{3}. [Illustration: Fig. 74] The mixture is heated below the fusing point for some time, so that theescaping carbon dioxide may not spatter the hot liquid; the heat is thenincreased and the mixture kept in a state of fusion until all gasesformed in the reaction have escaped. _Molding and blowing of glass. _ The way in which the melted mixture ishandled in the glass factory depends upon the character of the articleto be made. Many articles, such as bottles, are made by blowing theplastic glass into hollow molds of the desired shape. The mold is firstopened, as shown in Fig. 74. A lump of plastic glass A on the hollowrod B is lowered into the mold, which is then closed by the handlesC. By blowing into the tube the glass is blown into the shape of themold. The mold is then opened and the bottle lifted out. The neck of thebottle must be cut off at the proper place and the sharp edges roundedoff in a flame. Other objects, such as lamp chimneys, are made by getting a lump ofplastic glass on the end of a hollow iron rod and blowing it into thedesired shape without the help of a mold, great skill being required inthe manipulation of the glass. Window glass is made by blowing largehollow cylinders about 6 ft. Long and 1-1/2 ft. In diameter. These arecut longitudinally, and are then placed in an oven and heated until theysoften, when they are flattened out into plates (Fig. 75). Plate glassis cast into flat slabs, which are then ground and polished to perfectlyplane surfaces. _Varieties of glass. _ The ingredients mentioned above make a soft, easily fusible glass. If potassium carbonate is substituted for thesodium carbonate, the glass is much harder and less easily fused;increasing the amount of sand has somewhat the same effect. Potassiumglass is largely used in making chemical glassware, since it resists theaction of reagents better than the softer sodium glass. If lead oxide issubstituted for the whole or a part of the lime, the glass is very soft, but has a high index of refraction and is valuable for making opticalinstruments and artificial jewels. [Illustration: Fig. 75] _Coloring of glass. _ Various substances fused along with the glassmixture give characteristic colors. The amber color of common bottles isdue to iron compounds in the glass; in other cases iron colors the glassgreen. Cobalt compounds color it deep blue; those of manganese give itan amethyst tint and uranium compounds impart a peculiar yellowish greencolor. Since iron is nearly always present in the ingredients, glass isusually slightly yellow. This color can be removed by adding the properamount of manganese dioxide, for the amethyst color of manganese and theyellow of iron together produce white light. _Nature of glass. _ Glass is not a definite chemical compound and itscomposition varies between wide limits. Fused glass is really a solutionof various silicates, such as those of calcium and lead, in fused sodiumor potassium silicate. A certain amount of silicon dioxide is alsopresent. This solution is then allowed to solidify under such conditionsof cooling that the dissolved substances do not separate from thesolvent. The compounds which are used to color the glass are sometimesconverted into silicates, which then dissolve in the glass, giving it auniform color. In other cases, as in the milky glasses which resembleporcelain in appearance, the color or opaqueness is due to the finelydivided color material evenly distributed throughout the glass, but notdissolved in it. Milky glass is made by mixing calcium fluoride, tinoxide, or some other insoluble substance in the melted glass. Copper orgold in metallic form scattered through glass gives it shades of red. TITANIUM Titanium is a very widely distributed element in nature, being found in almost all soils, in many rocks, and even in plant and animal tissues. It is not very abundant in any one locality, and it possesses little commercial value save in connection with the iron industry. Its most common ore is rutile (TiO_{2}), which resembles silica in many respects. In both physical and chemical properties titanium resembles silicon, though it is somewhat more metallic in character. This resemblance is most marked in the acids of titanium. It not only forms metatitanic and orthotitanic acids but a great variety of polytitanic acids as well. BORON ~Occurrence. ~ Boron is never found free in nature. It occurs as boric acid(H_{3}BO_{3}), and in salts of polyboric acids, which usually have verycomplicated formulas. ~Preparation and properties. ~ Boron can be prepared from its oxide byreduction with magnesium, exactly as in the case of silicon. Itresembles silicon very strikingly in its properties. It occurs inseveral allotropic forms, is very hard when crystallized, and is ratherinactive toward reagents. It forms a hydride, BH_{3}, and combinesdirectly with the elements of the chlorine family. Boron fluoride(BF_{3}) is very similar to silicon fluoride in its mode of formationand chemical properties. ~Boric oxide~ (B_{2}O_{3}). Boron forms one well-known oxide, B_{2}O_{3}, called boric anhydride. It is formed as a glassy mass by heating boricacid to a high temperature. It absorbs water very readily, uniting withit to form boric acid again: B_{2}O_{3} + 3H_{2}O = 2H_{3}BO_{3}. In this respect it differs from silicon dioxide, which will not combinedirectly with water. ~Boric acid~ (H_{3}BO_{3}). This is found in nature in considerablequantities and forms one of the chief sources of boron compounds. It isfound dissolved in the water of hot springs in some localities, particularly in Italy. Being volatile with steam, the vapor whichescapes from these springs has some boric acid in it. It is easilyobtained from these sources by condensation and evaporation, thenecessary heat being supplied by other hot springs. Boric acid crystallizes in pearly flakes, which are greasy to the touch. In the laboratory it is easily prepared by treating a strong, hotsolution of borax with sulphuric acid. Boric acid being sparinglysoluble in water crystallizes out on cooling: Na_{2}B_{4}O_{7} + 5H_{2}O + H_{2}SO_{4} = Na_{2}SO_{4} + 4H_{3}BO_{3}. The substance is a mild antiseptic, and on this account is often used inmedicine and as a preservative for canned foods and milk. ~Metaboric and polyboric acids. ~ When boric acid is gently heated it isconverted into metaboric acid (HBO_{2}): H_{3}BO_{3} = HBO_{2} + H_{2}O. On heating metaboric acid to a somewhat higher temperature tetraboricacid (H_{2}B_{4}O_{7}) is formed: 4HBO_{2} = H_{2}B_{4}O_{7} + H_{2}O. Many other complex acids of boron are known. ~Borax. ~ Borax is the sodium salt of tetraboric acid, having the formulaNa_{2}B_{4}O_{7}·10 H_{2}O. It is found in some arid countries, assouthern California and Tibet, but is now made commercially from themineral colemanite, which is the calcium salt of a complex boric acid. When this is treated with a solution of sodium carbonate, calciumcarbonate is precipitated and borax crystallizes from the solution. When heated borax at first swells up greatly, owing to the expulsion ofthe water of crystallization, and then melts to a clear glass. Thisglass has the property of easily dissolving many metallic oxides, and onthis account borax is used as a flux in soldering, for the purpose ofremoving from the metallic surfaces to be soldered the film of oxidewith which they are likely to be covered. These oxides often give acharacteristic color to the clear borax glass, and borax beads aretherefore often used in testing for the presence of metals, instead ofthe metaphosphoric acid bead already described. The reason that metallic oxides dissolve in borax is that borax contains an excess of acid anhydride, as can be more easily seen if its formula is written 2NaBO_{2} + B_{2}O_{3}. The metallic oxide combines with this excess of acid anhydride, forming a mixed salt of metaboric acid. Borax is extensively used as a constituent of enamels and glazes forboth metal ware and pottery. It is also used as a flux in soldering andbrazing, and in domestic ways it serves as a mild alkali, as apreservative for meats, and in a great variety of less importantapplications. EXERCISES 1. Account for the fact that a solution of borax in water is alkaline. 2. What weight of water of crystallization does 1 kg. Of borax contain? 3. When a concentrated solution of borax acts on silver nitrate a borateof silver is formed. If the solution of borax is dilute, however, anhydroxide of silver forms. Account for this difference in behavior. CHAPTER XXII THE METALS ~The metals. ~ The elements which remain to be considered are knowncollectively as the metals. They are also called the base-formingelements, since their hydroxides are bases. A metal may therefore bedefined as an element whose hydroxide is a base. When a base dissolvesin water the hydroxyl groups form the anions, while the metallic elementforms the cations. From this standpoint a metal can be defined as anelement capable of forming simple cations in solution. The distinction between a metal and a non-metal is not a very sharp one, since the hydroxides of a number of elements act as bases under someconditions and as acids under others. We have seen that antimony is anelement of this kind. ~Occurrence of metals in nature. ~ A few of the metals are found in naturein the free state. Among these are gold, platinum, and frequentlycopper. They are usually found combined with other elements in the formof oxides or salts of various acids. Silicates, carbonates, sulphides, and sulphates are the most abundant salts. All inorganic substancesoccurring in nature, whether they contain a metal or not, are called_minerals_. Those minerals from which a useful substance can beextracted are called _ores_ of the substance. These two terms are mostfrequently used in connection with the metals. ~Extraction of metals, --metallurgy. ~ The process of extracting a metalfrom its ores is called the metallurgy of the metal. The metallurgy ofeach metal presents peculiarities of its own, but there are severalmethods of general application which are very frequently employed. 1. _Reduction of an oxide with carbon. _ Many of the metals occur innature in the form of oxides. When these oxides are heated to a hightemperature with carbon the oxygen combines with it and the metal is setfree. Iron, for example, occurs largely in the form of the oxideFe_{2}O_{3}. When this is heated with carbon the reaction expressed inthe following equation takes place: Fe_{2}O_{3} + 3 C = 2 Fe + 3 CO. Many ores other than oxides may be changed into oxides which can then bereduced by carbon. The conversion of such ores into oxides is generallyaccomplished by heating, and this process is called _roasting_. Manycarbonates and hydroxides decompose directly into the oxide on heating. Sulphides, on the other hand, must be heated in a current of air, theoxygen of the air entering into the reaction. The following equationswill serve to illustrate these changes in the case of the ores of iron: FeCO_{3} = FeO + CO_{2}, 2Fe(OH)_{3} = Fe_{2}O_{3} + 3H_{2}O, 2FeS_{2} + 11O = Fe_{2}O_{3} + 4SO_{2}. 2. _Reduction of an oxide with aluminium. _ Not all oxides, however, canbe reduced by carbon. In such cases aluminium may be used. Thus chromiummay be obtained in accordance with the following equation: Cr_{2}O_{3} + 2 Al = 2 Cr + Al_{2}O_{3}. This method is a comparatively new one, having been brought into use bythe German chemist Goldschmidt; hence it is sometimes called theGoldschmidt method. 3. _Electrolysis. _ In recent years increasing use is being made of theelectric current in the preparation of metals. In some cases theseparation of the metal from its compounds is accomplished by passingthe current through a solution of a suitable salt of the metal, themetal usually being deposited upon the cathode. In other cases thecurrent is passed through a fused salt of the metal, the chloride beingbest adapted to this purpose. ~Electro-chemical industries. ~ Most of the electro-chemical industries ofthe country are carried on where water power is abundant, since thisfurnishes the cheapest means for the generation of electrical energy. Niagara Falls is the most important locality in this country for suchindustries, and many different electro-chemical products aremanufactured there. Some industries depend upon electrolytic processes, while in others the electrical energy is used merely as a source of heatin electric furnaces. ~Preparation of compounds of the metals. ~ Since the compounds of themetals are so numerous and varied in character, there are many ways ofpreparing them. In many cases the properties of the substance to beprepared, or the material available for its preparation, suggest arather unusual way. There are, however, a number of general principleswhich are constantly applied in the preparation of the compounds of themetals, and a clear understanding of them will save much time and effortin remembering the details in any given case. The most important ofthese general methods for the preparation of compounds are thefollowing: 1. _By direct union of two elements. _ This is usually accomplished byheating the two elements together. Thus the sulphides, chlorides, andoxides of a metal can generally be obtained in this way. The followingequations serve as examples of this method: Fe + S = FeS, Mg + O = MgO, Cu + 2Cl = CuCl_{2}. 2. _By the decomposition of a compound. _ This decomposition may bebrought about either by heat alone or by the combined action of heat anda reducing agent. Thus when the nitrate of a metal is heated the oxideof the metal is usually obtained. Copper nitrate, for example, decomposes as follows: Cu(NO_{3})_{2} = CuO + 2NO_{2} + O. Similarly the carbonates of the metals yield oxides, thus: CaCO_{3} = CaO + CO_{2}. Most of the hydroxides form an oxide and water when heated: 2Al(OH)_{3} = Al_{2}O_{3} + 3H_{2}O. When heated with carbon, sulphates are reduced to sulphides, thus: BaSO_{4} + 2C = BaS + 2CO_{2}. 3. _Methods based on equilibrium in solution. _ In the preparation ofcompounds the first requisite is that the reactions chosen shall be ofsuch a kind as will go on to completion. In the chapter on chemicalequilibrium it was shown that reactions in solution may become completein either of three ways: (1) a gas may be formed which escapes fromsolution; (2) an insoluble solid may be formed which precipitates; (3)two different ions may combine to form undissociated molecules. By thejudicious selection of materials these principles may be applied to thepreparation of a great variety of compounds, and illustrations of suchmethods will very frequently be found in the subsequent pages. 4. _By fusion methods. _ It sometimes happens that substances which areinsoluble in water and in acids, and which cannot therefore be broughtinto double decomposition in the usual way, are soluble in otherliquids, and when dissolved in them can be decomposed and converted intoother desired compounds. Thus barium sulphate is not soluble in water, and sulphuric acid, being less volatile than most other acids, cannoteasily be driven out from this salt When brought into contact withmelted sodium carbonate, however, it dissolves in it, and since bariumcarbonate is insoluble in melted sodium carbonate, double decompositiontakes place: Na_{2}CO_{3} + BaSO_{4} = BaCO_{3} + Na_{2}SO_{4}. On dissolving the cooled mixture in water the sodium sulphate formed inthe reaction, together with any excess of sodium carbonate which may bepresent, dissolves. The barium carbonate can then be filtered off andconverted into any desired salt by the processes already described. 5. _By the action of metals on salts of other metals. _ When a strip ofzinc is placed in a solution of a copper salt the copper is precipitatedand an equivalent quantity of zinc passes into solution: Zn + CuSO_{4} = Cu + ZnSO_{4}. In like manner copper will precipitate silver from its salts: Cu + Ag_{2}SO_{4} = 2Ag + CuSO_{4}. It is possible to tabulate the metals in such a way that any one of themin the table will precipitate any one following it from its salts. Thefollowing is a list of some of the commoner metals arranged in this way: Zinc Iron Tin Lead Copper Bismuth Mercury Silver Gold According to this table copper will precipitate bismuth, mercury, silver, or gold from their salts, and will in turn be precipitated byzinc, iron, tin, or lead. Advantage is taken of this principle in thepurification of some of the metals, and occasionally in the preparationof metals and their compounds. ~Important insoluble compounds. ~ Since precipitates play so important apart in the reactions which substances undergo, as well as in thepreparation of many chemical compounds, it is important to know whatsubstances are insoluble. Knowing this, we can in many cases predictreactions under certain conditions, and are assisted in devising ways toprepare desired compounds. While there is no general rule which willenable one to foretell the solubility of any given compound, nevertheless a few general statements can be made which will be of muchassistance. 1. _Hydroxides. _ All hydroxides are insoluble save those of ammonium, sodium, potassium, calcium, barium, and strontium. 2. _Nitrates. _ All nitrates are soluble in water. 3. _Chlorides. _ All chlorides are soluble save silver and mercurouschlorides. (Lead chloride is but slightly soluble. ) 4. _Sulphates. _ All sulphates are soluble save those of barium, strontium, and lead. (Sulphates of silver and calcium are onlymoderately soluble. ) 5. _Sulphides. _ All sulphides are insoluble save those of ammonium, sodium, and potassium. The sulphides of calcium, barium, strontium, andmagnesium are insoluble in water, but are changed by hydrolysis intoacid sulphides which are soluble. On this account they cannot beprepared by precipitation. 6. _Carbonates, phosphates, and silicates. _ All normal carbonates, phosphates, and silicates are insoluble save those of ammonium, sodiumand potassium. EXERCISES 1. Write equations representing four different ways for preparingCu(NO_{3})_{2}. 2. Write equations representing six different ways for preparingZnSO_{4}. 3. Write equations for two reactions to illustrate each of the threeways in which reactions in solutions may become complete. 4. Give one or more methods for preparing each of the followingcompounds: CaCl_{2}, PbCl_{2}, BaSO_{4}, CaCO_{3}, (NH_{4})_{2}S, Ag_{2}S, PbO, Cu(OH)_{2} (for solubilities, see last paragraph ofchapter). State in each case the general principle involved in themethod of preparation chosen. CHAPTER XXIII THE ALKALI METALS ================================================================= | | | | | | SYMBOL | ATOMIC | DENSITY | MELTING | FIRST PREPARED | | WEIGHT | | POINT |__________|________|________|_________|_________|________________ | | | | |Lithium | Li | 7. 03 | 0. 59 | 186. ° | Davy 1820Sodium | Na | 23. 05 | 0. 97 | 97. 6° | " 1807Potassium | K | 39. 15 | 0. 87 | 62. 5° | " 1807Rubidium | Rb | 85. 5 | 1. 52 | 38. 5° | Bunsen 1861Cęsium | Cs | 132. 9 | 1. 88 | 26. 5° | " 1860================================================================= ~The family. ~ The metals listed in the above table constitute the evenfamily in Group I in the periodic arrangement of the elements, andtherefore form a natural family. The name alkali metals is commonlyapplied to the family for the reason that the hydroxides of the mostfamiliar members of the family, namely sodium and potassium, have longbeen called alkalis. 1. _Occurrence. _ While none of these metals occur free in nature, theircompounds are very widely distributed, being especially abundant in seaand mineral waters, in salt beds, and in many rocks. Only sodium andpotassium occur in abundance, the others being rarely found in anyconsiderable quantity. 2. _Preparation. _ The metals are most conveniently prepared by theelectrolysis of their fused hydroxides or chlorides, though it ispossible to prepare them by reducing their oxides or carbonates withcarbon. 3. _Properties. _ They are soft, light metals, having low melting pointsand small densities, as is indicated in the table. Their melting pointsvary inversely with their atomic weights, while their densities (sodiumexcepted) vary directly with these. The pure metals have a silveryluster but tarnish at once when exposed to the air, owing to theformation of a film of oxide upon the surface of the metal. They aretherefore preserved in some liquid, such as coal oil, which contains nooxygen. Because of their strong affinity for oxygen they decompose waterwith great ease, forming hydroxides and liberating hydrogen inaccordance with the equation M + H_{2}O = MOH + H, where M stands for any one of these metals. These hydroxides are whitesolids; they are readily soluble in water and possess very strong basicproperties. These bases are nearly equal in strength, that is, they alldissociate in water to about the same extent. 4. _Compounds. _ The alkali metals almost always act as univalentelements in the formation of compounds, the composition of which can berepresented by such formulas as MH, MCl, MNO_{3}, M_{2}SO_{4}, M_{3}PO_{4}. These compounds, when dissolved in water, dissociate insuch a way as to form simple, univalent metallic ions which arecolorless. With the exception of lithium these metals form very fewinsoluble compounds, so that it is not often that precipitatescontaining them are obtained. Only sodium and potassium will be studiedin detail, since the other metals of the family are of relatively smallimportance. The compounds of sodium and potassium are so similar in properties thatthey can be used interchangeably for most purposes. Other things beingequal, the sodium compounds are prepared in preference to those ofpotassium, since they are cheaper. When a given sodium compound isdeliquescent, or is so soluble that it is difficult to purify, thecorresponding potassium compound is prepared in its stead, provided itsproperties are more desirable in these respects. SODIUM ~Occurrence in nature. ~ Large deposits of sodium chloride have been foundin various parts of the world, and the water of the ocean and of manylakes and springs contains notable quantities of it. The element alsooccurs as a constituent of many rocks and is therefore present in thesoil formed by their disintegration. The mineral cryolite(Na_{3}AlF_{6}) is an important substance, and the nitrate, carbonate, and borate also occur in nature. ~Preparation. ~ In 1807 Sir Humphry Davy succeeded in preparing very smallquantities of metallic sodium by the electrolysis of the fusedhydroxide. On account of the cost of electrical energy it was for manyyears found more economical to prepare it by reducing the carbonate withcarbon in accordance with the following equation: Na_{2}CO_{3} + 2C = 2Na + 3CO. The cost of generating the electric current has been diminished to suchan extent, however, that it is now more economical to prepare sodium byDavy's original method, namely, by the electrolysis of the fusedhydroxide or chloride. When the chloride is used the process isdifficult to manage, owing to the higher temperature required to keepthe electrolyte fused, and because of the corroding action of the fusedchloride upon the containing vessel. [Illustration: SIR HUMPHRY DAVY (English) (1778-1829) Isolated sodium, lithium, potassium, barium, strontium, and calcium bymeans of electrolysis; demonstrated the elementary nature of chlorine;invented the safety lamp; discovered the stupefying effects of nitrousoxide] ~Technical preparation. ~ The sodium hydroxide is melted in a cylindrical iron vessel (Fig. 76) through the bottom of which rises the cathode K. The anodes A, several in number, are suspended around the cathode from above. A cylindrical vessel C floats in the fused alkali directly over the cathode, and under this cap the sodium and hydrogen liberated at the cathode collect. The hydrogen escapes by lifting the cover, and the sodium, protected from the air by the hydrogen, is skimmed or drained off from time to time. Oxygen is set free upon the anode and escapes into the air through the openings O without coming into contact with the sodium or hydrogen. This process is carried on extensively at Niagara Falls. [Illustration: Fig. 76] ~Properties. ~ Sodium is a silver-white metal about as heavy as water, andso soft that it can be molded easily by the fingers or pressed intowire. It is very active chemically, combining with most of thenon-metallic elements, such as oxygen and chlorine, with great energy. It will often withdraw these elements from combination with otherelements, and is thus able to decompose water and the oxides andchlorides of many metals. ~Sodium peroxide~ (NaO). Since sodium is a univalent element we shouldexpect it to form an oxide of the formula Na_{2}O. While such an oxidecan be prepared, the peroxide (NaO) is much better known. It is ayellowish-white powder made by burning sodium in air. Its chief use isas an oxidizing agent. When heated with oxidizable substances it givesup a part of its oxygen, as shown in the equation 2NaO = Na_{2}O + O. Water decomposes it in accordance with the equation 2NaO + 2H_{2}O = 2NaOH + H_{2}O_{2}. Acids act readily upon it, forming a sodium salt and hydrogen peroxide: 2NaO + 2HCl = 2NaCl + H_{2}O_{2}. In these last two reactions the hydrogen dioxide formed may decomposeinto water and oxygen if the temperature is allowed to rise: H_{2}O_{2} = H_{2}O + O. ~Peroxides. ~ It will be remembered that barium dioxide (BaO_{2}) yields hydrogen dioxide when treated with acids, and that manganese dioxide gives up oxygen when heated with sulphuric acid. Oxides which yield either hydrogen dioxide or oxygen when treated with water or an acid are called peroxides. ~Sodium hydroxide~ (_caustic soda_) (NaOH). 1. _Preparation. _ Sodiumhydroxide is prepared commercially by several processes. (a) In the older process, still in extensive use, sodium carbonate istreated with calcium hydroxide suspended in water. Calcium carbonate isprecipitated according to the equation Na_{2}CO_{3} + Ca(OH)_{2} = CaCO_{3} + 2NaOH. The dilute solution of sodium hydroxide, filtered from the calciumcarbonate, is evaporated to a paste and is then poured into molds tosolidify. It is sold in the form of slender sticks. (b) The newer methods depend upon the electrolysis of sodium chloride. In the Castner process a solution of salt is electrolyzed, the reactionbeing expressed as follows: NaCl + H_{2}O = NaOH + H + Cl. The chlorine escapes as a gas, and by an ingenious mechanical device thesodium hydroxide is prevented from mixing with the salt in the solution. In the Acker process the electrolyte is _fused_ sodium chloride. Thechlorine is evolved as a gas at the anode, while the sodium alloys withthe melted lead which forms the cathode. When this alloy is treated withwater the following reaction takes place: Na + H_{2}O = NaOH + H. [Illustration: Fig. 77] ~Technical process. ~ A sketch of an Acker furnace is represented in Fig. 77. The furnace is an irregularly shaped cast-iron box, divided intothree compartments, A, B, and C. Compartment A is lined withmagnesia brick. Compartments B and C are filled with melted lead, which also covers the bottom of A to a depth of about an inch. Abovethis layer in A is fused salt, into which dip carbon anodes D. Themetallic box and melted lead is the cathode. When the furnace is in operation chlorine is evolved at the anodes, and is drawn away through a pipe (not represented) to the bleaching-powder chambers. Sodium is set free at the surface of the melted lead in A, and at once alloys with it. Through the pipe E a powerful jet of steam is driven through the lead in B upwards into the narrow tube F. This forces the lead alloy up through the tube and over into the chamber G. In this process the steam is decomposed by the sodium in the alloy, forming melted sodium hydroxide and hydrogen. The melted lead and sodium hydroxide separate into two layers in G, and the sodium hydroxide, being on top, overflows into tanks from which it is drawn off and packed in metallic drums. The lead is returned to the other compartments of the furnace by a pipe leading from H to I. Compartment C serves merely as a reservoir for excess of melted lead. 2. _Properties. _ Sodium hydroxide is a white, crystalline, brittlesubstance which rapidly absorbs water and carbon dioxide from the air. As the name (caustic soda) indicates, it is a very corrosive substance, having a disintegrating action on most animal and vegetable tissues. Itis a strong base. It is used in a great many chemical industries, andunder the name of lye is employed to a small extent as a cleansing agentfor household purposes. ~Sodium chloride~ (_common salt_) (NaCl). 1. _Preparation. _ Sodiumchloride, or common salt, is very widely distributed in nature. Thickstrata, evidently deposited at one time by the evaporation of saltwater, are found in many places. In the United States the most importantlocalities for salt are New York, Michigan, Ohio, and Kansas. Sometimesthe salt is mined, especially if it is in the pure form called rocksalt. More frequently a strong brine is pumped from deep wells sunk intothe salt deposit, and is then evaporated in large pans until the saltcrystallizes out. The crystals are in the form of small cubes andcontain no water of crystallization; some water is, however, held incavities in the crystals and causes the salt to decrepitate when heated. 2. _Uses. _ Since salt is so abundant in nature it forms the startingpoint in the preparation of all compounds containing either sodium orchlorine. This includes many substances of the highest importance tocivilization, such as soap, glass, hydrochloric acid, soda, andbleaching powder. Enormous quantities of salt are therefore producedeach year. Small quantities are essential to the life of man andanimals. Pure salt does not absorb moisture; the fact that ordinary saltbecomes moist in air is not due to a property of the salt, but toimpurities commonly occurring in it, especially calcium and magnesiumchlorides. ~Sodium sulphate~ (_Glauber's salt_) (Na_{2}SO_{4}·10H_{2}O). This salt isprepared by the action of sulphuric acid upon sodium chloride, hydrochloric acid being formed at the same time: 2NaCl + H_{2}SO_{4} = Na_{2}SO_{4} + 2HCl. Some sodium sulphate is prepared by the reaction represented in theequation MgSO_{4} + 2NaCl = Na_{2}SO_{4} + MgCl_{2}. The magnesium sulphate required for this reaction is obtained in largequantities in the manufacture of potassium chloride, and being of littlevalue for any other purpose is used in this way. The reaction dependsupon the fact that sodium sulphate is the least soluble of any of thefour factors in the equation, and therefore crystallizes out when hot, saturated solutions of magnesium sulphate and sodium chloride are mixedtogether and the resulting mixture cooled. Sodium sulphate forms large efflorescent crystals. The salt isextensively used in the manufacture of sodium carbonate and glass. Smallquantities are used in medicine. ~Sodium sulphite~ (Na_{2}SO_{3}·7H_{2}O). Sodium sulphite is prepared bythe action of sulphur dioxide upon solutions of sodium hydroxide, thereaction being analogous to the action of carbon dioxide upon sodiumhydroxide. Like the carbonate, the sulphite is readily decomposed byacids: Na_{2}SO_{3} + 2HCl = 2NaCl + H_{2}O + SO_{2}. Because of this reaction sodium sulphite is used as a convenient sourceof sulphur dioxide. It is also used as a disinfectant and apreservative. ~Sodium thiosulphate~ (_hyposulphite of soda or "hypo"_)(Na_{2}S_{2}O_{3}·5H_{2}O). This salt, commonly called sodiumhyposulphite, or merely hypo, is made by boiling a solution of sodiumsulphite with sulphur: Na_{2}SO_{3} + S = Na_{2}S_{2}O_{3}. It is used in photography and in the bleaching industry, to absorb theexcess of chlorine which is left upon the bleached fabrics. ~Thio compounds. ~ The prefix "thio" means sulphur. It is used to designate substances which may be regarded as derived from oxygen compounds by replacing the whole or a part of their oxygen with sulphur. The thiosulphates may be regarded as sulphates in which one atom of oxygen has been replaced by an atom of sulphur. This may be seen by comparing the formula Na_{2}SO_{4} (sodium sulphate) with the formula Na_{2}S_{2}O_{3} (sodium thiosulphate). ~Sodium carbonate~ (_sal soda_)(Na_{2}CO_{3}·10H_{2}O). There are twodifferent methods now employed in the manufacture of this importantsubstance. 1. _Le Blanc process. _ This older process involves several distinctreactions, as shown in the following equations. (a) Sodium chloride is first converted into sodium sulphate: 2NaCl + H_{2}SO_{4} = Na_{2}SO_{4} + 2HCl. (b) The sodium sulphate is next reduced to sulphide by heating it withcarbon: Na_{2}SO_{4} + 2C = Na_{2}S + 2CO_{2}. (c) The sodium sulphide is then heated with calcium carbonate, whendouble decomposition takes place: Na_{2}S + CaCO_{3} = CaS + Na_{2}CO_{3}. ~Technical preparation of sodium carbonate. ~ In a manufacturing plant the last two reactions take place in one process. Sodium sulphate, coal, and powdered limestone are heated together to a rather high temperature. The coal reduces the sulphate to sulphide, which in turn reacts upon the calcium carbonate. Some limestone is decomposed by the heat, forming calcium oxide. When treated with water the calcium oxide is changed into hydroxide, and this prevents the water from decomposing the insoluble calcium sulphide. The crude product of the process is a hard black cake called black ash. On digesting this mass with water the sodium carbonate passes into solution. The pure carbonate is obtained by evaporation of this solution, crystallizing from it in crystals of the formula Na_{2}CO_{3}·10H_{2}O. Since over 60% of this salt is water, the crystals are sometimes heated until it is driven off. The product is called calcined soda, and is, of course, more valuable than the crystallized salt. 2. _Solvay process. _ This more modern process depends upon the reactionsrepresented in the equations NaCl + NH_{4}HCO_{3} = NaHCO_{3} + NH_{4}Cl, 2NaHCO_{3} = Na_{2}CO_{3} + H_{2}O + CO_{2}. The reason the first reaction takes place is that sodium hydrogencarbonate is sparingly soluble in water, while the other compounds arefreely soluble. When strong solutions of sodium chloride and of ammoniumhydrogen carbonate are brought together the sparingly soluble sodiumhydrogen carbonate is precipitated. This is converted into the normalcarbonate by heating, the reaction being represented in the secondequation. ~Technical preparation. ~ In the Solvay process a very concentrated solution of salt is first saturated with ammonia gas, and a current of carbon dioxide is then conducted into the solution. In this way ammonium hydrogen carbonate is formed: NH_{3} + H_{2}O + CO_{2} = NH_{4}HCO_{3}. This enters into double decomposition with the salt, as shown in the first equation under the Solvay process. After the sodium hydrogen carbonate has been precipitated the mother liquors containing ammonium chloride are treated with lime: 2NH_{4}Cl + CaO = CaCl_{2} + 2 NH_{3} + H_{2}O. The lime is obtained by burning limestone: CaCO_{3} = CaO + CO_{2}. The ammonia and carbon dioxide evolved in the latter two reactions are used in the preparation of an additional quantity of ammonium hydrogen carbonate. It will thus be seen that there is no loss of ammonia. The only materials permanently used up are calcium carbonate and salt, while the only waste product is calcium chloride. ~Historical. ~ In former times sodium carbonate was made by burning seaweeds and extracting the carbonate from their ash. On this account the salt was called _soda ash_, and the name is still in common use. During the French Revolution this supply was cut off, and in behalf of the French government Le Blanc made a study of methods of preparing the carbonate directly from salt. As a result he devised the method which bears his name, and which was used exclusively for many years. It has been replaced to a large extent by the Solvay process, which has the advantage that the materials used are inexpensive, and that the ammonium hydrogen carbonate used can be regenerated from the products formed in the process. Much expense is also saved in fuel, and the sodium hydrogen carbonate, which is the first product of the process, has itself many commercial uses. The Le Blanc process is still used, however, since the hydrochloric acid generated is of value. ~By-products. ~ The substances obtained in a given process, aside from the main product, are called the by-products. The success of many processes depends upon the value of the by-products formed. Thus hydrochloric acid, a by-product in the Le Blanc process, is valuable enough to make the process pay, even though sodium carbonate can be made cheaper in other ways. ~Properties of sodium carbonate. ~ Sodium carbonate forms large crystals ofthe formula Na_{2}CO_{3} · 10 H_{2}O. It has a mild alkaline reactionand is used for laundry purposes under the name of washing soda. Meremention of the fact that it is used in the manufacture of glass, soap, and many chemical reagents will indicate its importance in theindustries. It is one of the few soluble carbonates. ~Sodium hydrogen carbonate~ (_bicarbonate of soda_) (NaHCO_{3}). Thissalt, commonly called bicarbonate of soda, or baking soda, is made bythe Solvay process, as explained above, or by passing carbon dioxideinto strong solutions of sodium carbonate: Na_{2}CO_{3} + H_{2}O + CO_{2} = 2NaHCO_{3}. The bicarbonate, being sparingly soluble, crystallizes out. A mixture ofthe bicarbonate with some substance (the compound known as cream oftartar is generally used) which slowly reacts with it, liberating carbondioxide, is used largely in baking. The carbon dioxide generated forcesits way through the dough, thus making it porous and light. ~Sodium nitrate~ (_Chili saltpeter_) (NaNO_{3}). This substance is foundin nature in arid regions in a number of places, where it has beenformed apparently by the decay of organic substances in the presence ofair and sodium salts. The largest deposits are in Chili, and most of thenitrate of commerce comes from that country. Smaller deposits occur inCalifornia and Nevada. The commercial salt is prepared by dissolving thecrude nitrate in water, allowing the insoluble earthy materials tosettle, and evaporating the clear solution so obtained tocrystallization. The soluble impurities remain for the most part in themother liquors. Since this salt is the only nitrate found extensively in nature, it isthe material from which other nitrates as well as nitric acid areprepared. It is used in enormous quantities in the manufacture ofsulphuric acid and potassium nitrate, and as a fertilizer. ~Sodium phosphate~ (Na_{2}HPO_{4}·12H_{2}O). Since phosphoric acid hasthree replaceable hydrogen atoms, three sodium phosphates arepossible, --two acid salts and one normal. All three can be made withoutdifficulty, but disodium phosphate is the only one which is largelyused, and is the salt which is commonly called sodium phosphate. It ismade by the action of phosphoric acid on sodium carbonate: Na_{2}CO_{3} + H_{3}PO_{4} = Na_{2}HPO_{4} + CO_{2} + H_{2}O. It is interesting as being one of the few phosphates which are solublein water, and is the salt commonly used when a soluble phosphate isneeded. ~Normal sodium phosphate~ (Na_{3}PO_{4}). Although this is a normal saltits solution has a strongly alkaline reaction. This is due to the factthat the salt hydrolyzes in solution into sodium hydroxide and disodiumphosphate, as represented in the equation Na_{3}PO_{4} + H_{2}O = Na_{2}HPO_{4} + NaOH. Sodium hydroxide is strongly alkaline, while disodium phosphate isnearly neutral in reaction. The solution as a whole is thereforealkaline. The salt is prepared by adding a large excess of sodiumhydroxide to a solution of disodium phosphate and evaporating tocrystallization. The excess of the sodium hydroxide reverses thereaction of hydrolysis and the normal salt crystallizes out. ~Sodium tetraborate ~(_borax_) (Na_{2}B_{4}O_{7}·10H_{2}O). The propertiesof this important compound have been discussed under the head of boron. POTASSIUM ~Occurrence in nature. ~ Potassium is a constituent of many common rocksand minerals, and is therefore a rather abundant element, though not soabundant as sodium. Feldspar, which occurs both by itself and as aconstituent of granite, contains considerable potassium. The element isa constituent of all clay and of mica and also occurs in very largedeposits at Stassfurt, Germany, in the form of the chloride andsulphate, associated with compounds of sodium and magnesium. In smallquantities it is found as nitrate and in many other forms. The natural decomposition of rocks containing potassium gives rise tovarious compounds of the element in all fertile soils. Its solublecompounds are absorbed by growing plants and built up into complexvegetable substances; when these are burned the potassium remains in theash in the form of the carbonate. Crude carbonate obtained from woodashes was formerly the chief source of potassium compounds; they are nowmostly prepared from the salts of the Stassfurt deposits. ~Stassfurt salts. ~ These salts form very extensive deposits in middle and north Germany, the most noted locality for working them being at Stassfurt. The deposits are very thick and rest upon an enormous layer of common salt. They are in the form of a series of strata, each consisting largely of a single mineral salt. A cross section of these deposits is shown in Fig. 78. While these strata are salts from a chemical standpoint, they are as solid and hard as many kinds of stone, and are mined as stone or coal would be. Since the strata differ in general appearance, each can be mined separately, and the various minerals can be worked up by methods adapted to each particular case. The chief minerals of commercial importance in these deposits are the following: Sylvine KCl. Anhydrite CaSO_{4}. Carnallite KCl·MgCl_{2}·6H_{2}O. Kainite K_{2}SO_{4}·MgSO_{4}·MgCl_{2}·6H_{2}O. Polyhalite K_{2}SO_{4}·MgSO_{4}·2CaSO_{4}·2H_{2}O. Kieserite MgSO_{4}·H_{2}O. Schönite K_{2}SO_{4}·MgSO_{4}·6H_{2}O. ~Preparation and properties. ~ The metal is prepared by the same methodused in the preparation of sodium. In most respects it is very similarto sodium, the chief difference being that it is even more energetic inits action upon other substances. The freshly cut, bright surfaceinstantly becomes dim through oxidation by the air. It decomposes watervery vigorously, the heat of reaction being sufficient to ignite thehydrogen evolved. It is somewhat lighter than sodium and is preservedunder gasoline. [Illustration: Fig. 78] ~Potassium hydroxide~ (_caustic potash_) (KOH). Potassium hydroxide isprepared by methods exactly similar to those used in the preparation ofsodium hydroxide, which compound it closely resembles in both physicaland chemical properties. It is not used to any very great extent, beingreplaced by the cheaper sodium hydroxide. ~Action of the halogen elements on potassium hydroxide. ~ When any one ofthe three halogen elements--chlorine, bromine, and iodine--is added to asolution of potassium hydroxide a reaction takes place, the nature ofwhich depends upon the conditions of the experiment. Thus, when chlorineis passed into a cold dilute solution of potassium hydroxide thereaction expressed by the following equation takes place: (1) 2KOH + 2Cl = KCl + KClO + H_{2}O. If the solution of hydroxide is concentrated and hot, on the other hand, the potassium hypochlorite formed according to equation (1) breaks downas fast as formed: (2) 3KClO = KClO_{3} + 2KCl. Equation (1), after being multiplied by 3, may be combined with equation(2), giving the following: (3) 6KOH + 6Cl = 5KCl + KClO_{3} + 3H_{2}O. This represents in a single equation the action of chlorine on hot, concentrated solutions of potassium hydroxide. By means of thesereactions one can prepare potassium chloride, potassium hypochlorite, and potassium chlorate. By substituting bromine or iodine for chlorinethe corresponding compounds of these elements are obtained. Some ofthese compounds can be obtained in cheaper ways. If the halogen element is added to a solution of sodium hydroxide orcalcium hydroxide, the reaction which takes place is exactly similar tothat which takes place with potassium hydroxide. It is possible, therefore, to prepare in this way the sodium and calcium compoundscorresponding to the potassium compounds given above. ~Potassium chloride~ (KCl). This salt occurs in nature in sea water, inthe mineral sylvine, and, combined with magnesium chloride, ascarnallite (KCl·MgCl_{2}·6H_{2}O). It is prepared from carnallite bysaturating boiling water with the mineral and allowing the solution tocool. The mineral decomposes while in solution, and the potassiumchloride crystallizes out on cooling, while the very soluble magnesiumchloride remains in solution. The salt is very similar to sodiumchloride both in physical and chemical properties. It is used in thepreparation of nearly all other potassium salts, and, together withpotassium sulphate, is used as a fertilizer. ~Potassium bromide~ (KBr). When bromine is added to a hot concentratedsolution of potassium hydroxide there is formed a mixture of potassiumbromide and potassium bromate in accordance with the reactions alreadydiscussed. There is no special use for the bromate, so the solution isevaporated to dryness, and the residue, consisting of a mixture of thebromate and bromide, is strongly heated. This changes the bromate tobromide, as follows: KBrO_{3} = KBr +3O. The bromide is then crystallized from water, forming large colorlesscrystals. It is used in medicine and in photography. ~Potassium iodide~ (KI). Potassium iodide may be made by exactly the samemethod as has just been described for the bromide, substituting iodinefor bromine. It is more frequently made as follows. Iron filings aretreated with iodine, forming the compound Fe_{3}I_{8}; on boiling thissubstance with potassium carbonate the reaction represented in thefollowing equation occurs: Fe_{3}I_{8} + 4K_{2}CO_{3} = Fe_{3}O_{4} + 8KI + 4CO_{2}. Potassium iodide finds its chief use in medicine. ~Potassium chlorate~ (KClO_{3}). This salt, as has just been explained, can be made by the action of chlorine on strong potassium hydroxidesolutions. The chief use of potassium chlorate is as an oxidizing agentin the manufacture of matches, fireworks, and explosives; it is alsoused in the preparation of oxygen and in medicine. ~Commercial preparation. ~ By referring to the reaction between chlorine and hot concentrated solutions of potassium hydroxide, it will be seen that only one molecule of potassium chlorate is formed from six molecules of potassium hydroxide. Partly because of this poor yield and partly because the potassium hydroxide is rather expensive, this process is not an economical one for the preparation of potassium chlorate. The commercial method is the following. Chlorine is passed into hot solutions of calcium hydroxide, a compound which is very cheap. The resulting calcium chloride and chlorate are both very soluble. To the solution of these salts potassium chloride is added, and as the solution cools the sparingly soluble potassium chlorate crystallizes out: Ca(ClO_{3})_{2} + 2KCl = 2KClO_{3} + CaCl_{2}. Electro-chemical processes are also used. ~Potassium nitrate~ (_saltpeter_) (KNO_{3}). This salt was formerly madeby allowing animal refuse to decompose in the open air in the presenceof wood ashes or earthy materials containing potassium. Under theseconditions the nitrogen in the organic matter is in part converted intopotassium nitrate, which was obtained by extracting the mass with waterand evaporating to crystallization. This crude and slow process is nowalmost entirely replaced by a manufacturing process in which thepotassium salt is made from Chili saltpeter: NaNO_{3} + KCl = NaCl + KNO_{3}. This process has been made possible by the discovery of the Chili niterbeds and the potassium chloride of the Stassfurt deposits. The reaction depends for its success upon the apparently insignificant fact that sodium chloride is almost equally soluble in cold and hot water. All four factors in the equation are rather soluble in cold water, but in hot water sodium chloride is far less soluble than the other three. When hot saturated solutions of sodium nitrate and potassium chloride are brought together, sodium chloride precipitates and can be filtered off, leaving potassium nitrate in solution, together with some sodium chloride. On cooling, potassium nitrate crystallizes out, leaving small amounts of the other salts in solution. Potassium nitrate is a colorless salt which forms very large crystals. It is stable in the air, and when heated is a good oxidizing agent, giving up oxygen quite readily. Its chief use is in the manufacture ofgunpowder. ~Gunpowder. ~ The object sought for in the preparation of gunpowder is to secure a solid substance which will remain unchanged under ordinary conditions, but which will explode readily when ignited, evolving a large volume of gas. When a mixture of carbon and potassium nitrate is ignited a great deal of gas is formed, as will be seen from the equation 2KNO_{3} + 3C = CO_{2} + CO + N_{2} + K_{2}CO_{3}. By adding sulphur to the mixture the volume of gas formed in the explosion is considerably increased: 2KNO_{3} + 3C + S = 3CO_{2} + N_{2} + K_{2}S. Gunpowder is simply a mechanical mixture of these three substances in the proportion required for the above reaction. While the equation represents the principal reaction, other reactions also take place. The gases formed in the explosion, when measured under standard conditions, occupy about two hundred and eighty times the volume of the original powder. Potassium sulphide (K_{2}S) is a solid substance, and it is largely due to it that gunpowder gives off smoke and soot when it explodes. Smokeless powder consists of organic substances which, on explosion, give only colorless gases, and hence produce no smoke. Sodium nitrate is cheaper than potassium nitrate, but it is not adapted to the manufacture of the best grades of powder, since it is somewhat deliquescent and does not give up its oxygen so readily as does potassium nitrate. It is used, however, in the cheaper grades of powder, such as are employed for blasting. ~Potassium cyanide~ (KCN). When animal matter containing nitrogen isheated with iron and potassium carbonate, complicated changes occurwhich result in the formation of a substance commonly called yellowprussiate of potash, which has the formula K_{4}FeC_{6}N_{6}. When thissubstance is heated with potassium, potassium cyanide is formed: K_{4}FeC_{6}N_{6} + 2 K = 6KCN + Fe. Since sodium is much cheaper than potassium it is often used in place ofit: K_{4}FeC_{6}N_{6} + 2Na = 4KCN + 2NaCN + Fe. The mixture of cyanides so resulting serves most of the purposes of thepure salt. It is used very extensively in several metallurgicalprocesses, particularly in the extraction of gold. Potassium cyanide isa white solid characterized by its poisonous properties, and must beused with extreme caution. ~Potassium carbonate~ (_potash_) (K_{2}CO_{3}). This compound occurs inwood ashes in small quantities. It cannot be prepared by the Solvayprocess, since the acid carbonate is quite soluble in water, but is madeby the Le Blanc process. Its chief use is in the manufacture of otherpotassium salts. ~Other salts of potassium. ~ Among the other salts of potassium frequentlymet with are the sulphate (K_{2}SO_{4}), the acid carbonate (KHCO_{3}), the acid sulphate (KHSO_{4}), and the acid sulphite (KHSO_{3}). Theseare all white solids. LITHIUM, RUBIDIUM, CĘSIUM Of the three remaining elements of the family--lithium, rubidium, andcęsium--lithium is by far the most common, the other two being veryrare. Lithium chloride and carbonate are not infrequently found innatural mineral waters, and as these substances are supposed to increasethe medicinal value of the water, they are very often added toartificial mineral waters in small quantities. COMPOUNDS OF AMMONIUM ~General. ~ As explained in a previous chapter, when ammonia is passed intowater the two compounds combine to form the base NH_{4}OH, known asammonium hydroxide. When this base is neutralized with acids there areformed the corresponding salts, known as the ammonium salts. Since theammonium group is univalent, ammonium salts resemble those of the alkalimetals in formulas; they also resemble the latter salts very much intheir chemical properties, and may be conveniently described inconnection with them. Among the ammonium salts the chloride, sulphate, carbonate, and sulphide are the most familiar. ~Ammonium chloride~ (_sal ammoniac_) (NH_{4}Cl). This substance isobtained by neutralizing ammonium hydroxide with hydrochloric acid. Itis a colorless substance crystallizing in fine needles, and, like mostammonium salts, is very soluble in water. When placed in a tube andheated strongly it decomposes into hydrochloric acid and ammonia. Whenthese gases reach a cooler portion of the tube they at once recombine, and the resulting ammonium chloride is deposited on the sides of thetube. In this way the salt can be separated from nonvolatile impurities. Ammonium chloride is sometimes used in preparation of ammonia; it isalso used in making dry batteries and in the laboratory as a chemicalreagent. ~Ammonium sulphate~ ((NH_{4})_{2}SO_{4}). This salt resembles the chloridevery closely, and, being cheaper, is used in place of it when possible. It is used in large quantity as a fertilizer, the nitrogen which itcontains being a very valuable food for plants. ~Ammonium carbonate~ ((NH_{4})_{2}CO_{3}). This salt, as well as the acidcarbonate (NH_{4}HCO_{3}), is used as a chemical reagent. They arecolorless solids, freely soluble in water. The normal carbonate is madeby heating ammonium chloride with powdered limestone (calciumcarbonate), the ammonium carbonate being obtained as a sublimate incompact hard masses: 2NH_{4}Cl + CaCO_{3} = (NH_{4})_{2}CO_{3} + CaCl_{2}. The salt always smells of ammonia, since it slowly decomposes, as shownin the equation (NH_{4})_{2}CO_{3} = NH_{4}HCO_{3} + NH_{3}. The acid carbonate, or bicarbonate, is prepared by saturating a solutionof ammonium hydroxide with carbon dioxide: NH_{4}OH + CO_{2} = NH_{4}HCO_{3}. It is a well-crystallized stable substance. ~Ammonium sulphide~ ((NH_{4})_{2}S). Ammonium sulphide is prepared by theaction of hydrosulphuric acid upon ammonium hydroxide: 2NH_{4}OH + H_{2}S = (NH_{4})_{2}S + 2H_{2}O. If the action is allowed to continue until no more hydrosulphuric acidis absorbed, the product is the acid sulphide, sometimes called thehydrosulphide: NH_{4}OH + H_{2}S = NH_{4}HS + H_{2}O. If equal amounts of ammonium hydroxide and ammonium acid sulphide arebrought together, the normal sulphide is formed: NH_{4}OH + NH_{4}HS = (NH_{4})_{2}S + H_{2}O It has been obtained in the solid state, but only with great difficulty. As used in the laboratory it is always in the form of a solution. It ismuch used in the process of chemical analysis because it is a solublesulphide and easily prepared. On exposure to the air ammonium sulphideslowly decomposes, being converted into ammonia, water, and sulphur: (NH_{4})_{2}S + O = 2NH_{3} + H_{2}O + S. As fast as the sulphur is liberated it combines with the unchangedsulphide to form several different ammonium sulphides in which there arefrom two to five sulphur atoms in the molecule, thus: (NH_{4})_{2}S_{2}, (NH_{4})_{2}S_{3}, (NH_{4})_{2}S_{5}. These sulphides in turn decomposeby further action of oxygen, so that the final products of the reactionare those given in the equation. A solution of these compounds is yellowand is sometimes called _yellow ammonium sulphide_. FLAME REACTION--SPECTROSCOPE When compounds of either sodium or potassium are brought into the non-luminous flame of a Bunsen burner the flame becomes colored. Sodium compounds color it intensely yellow, while those of potassium color it pale violet. When only one of these elements is present it is easy to identify it by this simple test, but when both are present the intense color of the sodium flame entirely conceals the pale tint characteristic of potassium compounds. It is possible to detect the potassium flame in such cases, however, in the following way. When light is allowed to shine through a very small hole or slit in some kind of a screen, such as a piece of metal, upon a triangular prism of glass, the light is bent or refracted out of its course instead of passing straight through the glass. It thus comes out of the prism at some angle to the line at which it entered. Yellow light is bent more than red, and violet more than yellow. When light made up of the yellow of sodium and the violet of potassium shines through a slit upon such a prism, the yellow and the violet lights come out at somewhat different angles, and so two colored lines of light--a yellow line and a violet line--are seen on looking into the prism in the proper direction. The instrument used for separating the rays of light in this way is called a _spectroscope_ (Fig. 79). The material to be tested is placed on a platinum wire and held in the colorless Bunsen flame. The resulting light passes through the slit in the end of tube B, and then through B to the prism. The resulting lines of light are seen by looking into the tube A, which contains a magnifying lens. Most elements give more than one image of the slit, each having a different color, and the series of colored lines due to an element is called its spectrum. [Illustration: Fig. 79] The spectra of the known elements have been carefully studied, and anyelement which imparts a characteristic color to a flame, or has aspectrum of its own, can be identified even when other elements arepresent. Through the spectroscopic examination of certain minerals anumber of elements have been discovered by the observation of lineswhich did not belong to any known element. A study of the substance thenbrought to light the new element. Rubidium and cęsium were discovered inthis way, rubidium having bright red lines and cęsium a very intenseblue line. Lithium colors the flame deep red, and has a bright red linein its spectrum. EXERCISES 1. What is an alkali? Can a metal itself be an alkali? 2. Write equations showing how the following changes may be broughtabout, giving the general principle involved in each change: NaCl -->Na_{2}SO_{3}, Na_{2}SO_{3} --> NaCl, NaCl --> NaBr, Na_{2}SO_{4} -->NaNO_{3}, NaNO_{3} --> NaHCO_{3}. 3. What carbonates are soluble? 4. State the conditions under which the reaction represented by thefollowing equation can be made to go in either direction: Na_{2}CO_{3} + H_{2}O + CO_{2} 2 NaHCO_{3}. 5. Account for the fact that solutions of sodium carbonate and potassiumcarbonate are alkaline. 6. What non-metallic element is obtained from the deposits of Chilisaltpeter? 7. Supposing concentrated hydrochloric acid (den. = 1. 2) to be worth sixcents a pound, what is the value of the acid generated in thepreparation of 1 ton of sodium carbonate by the Le Blanc process? 8. What weight of sodium carbonate crystals will 1 kg. Of the anhydroussalt yield? 9. Write equations for the preparation of potassium hydroxide by threedifferent methods. 10. What would take place if a bit of potassium hydroxide were leftexposed to the air? 11. Write the equations for the reactions between sodium hydroxide andbromine; between potassium hydroxide and iodine. 12. Write equations for the preparation of potassium sulphate; ofpotassium acid carbonate. 13. What weight of carnallite would be necessary in the preparation of 1ton of potassium carbonate? 14. Write the equations showing how ammonium chloride, ammoniumsulphate, ammonium carbonate, and ammonium nitrate may be prepared fromammonium hydroxide. 15. Write an equation to represent the reaction involved in thepreparation of ammonia from ammonium chloride. 16. What substances already studied are prepared from the followingcompounds? ammonium chloride; ammonium nitrate; ammonium nitrite; sodiumnitrate; sodium chloride. 17. How could you prove that the water in crystals of common salt is notwater of crystallization? 18. How could you distinguish between potassium chloride and potassiumiodide? between sodium chloride and ammonium chloride? between sodiumnitrate and potassium nitrate? [Illustration: ROBERT WILHELM BUNSEN (German) (1811-1899) Invented many lecture-room and laboratory appliances (Bunsen burner);invented the spectroscope and with it discovered rubidium and cęsium;greatly perfected methods of electrolysis, inventing a new battery; mademany investigations among metallic and organic substances] CHAPTER XXIV THE ALKALINE-EARTH FAMILY =========================================================================== | | | | | | | | | MILLIGRAMS SOL- | | | | | UBLE IN 1 L. | | | | | OF WATER AT 18° | | SYMBOL | ATOMIC | DENSITY |__________________| CARBONATE | | WEIGHT | | | | DECOMPOSES | | | | SULPHATE| HYDROX-| | | | | | IDE |__________|________|________|_________|_________|________|_________________ | | | | | |Calcium | Ca | 40. 1 | 1. 54 | 2070. 00 | 1670. | At dull red heatStrontium | Sr | 87. 6 | 2. 50 | 170. 00 | 7460. | At white heatBarium | Ba | 137. 4 | 3. 75 | 2. 29 | 36300. | Scarcely at all=========================================================================== ~The family. ~ The alkaline-earth family consists of the very abundantelement calcium and the much rarer elements strontium and barium. Theyare called the alkaline-earth metals because their properties arebetween those of the alkali metals and the earth metals. The earthmetals will be discussed in a later chapter. The family is alsofrequently called the calcium family. 1. _Occurrence. _ These elements do not occur free in nature. Their mostabundant compounds are the carbonates and sulphates; calcium also occursin large quantities as the phosphate and silicate. 2. _Preparation. _ The metals were first prepared by Davy in 1808 byelectrolysis. This method has again come into use in recent years. Strontium and barium have as yet been obtained only in small quantitiesand in the impure state, and many of their physical properties, such astheir densities and melting points, are therefore imperfectly known. 3. _Properties. _ The three metals resemble each other very closely. Theyare silvery-white in color and are about as hard as lead. Theirdensities increase with their atomic weights, as is shown in the tableon opposite page. Like the alkali metals they have a strong affinity foroxygen, tarnishing in the air through oxidation. They decompose water atordinary temperatures, forming hydroxides and liberating hydrogen. Whenignited in the air they burn with brilliancy, forming oxides of thegeneral formula MO. These oxides readily combine with water, accordingto the equation MO + H_{2}O = M(OH)_{2}. Each of the elements has a characteristic spectrum, and the presence ofthe metals can easily be detected by the spectroscope. 4. _Compounds. _ The elements are divalent in almost all of theircompounds, and these compounds in solution give simple, divalent, colorless ions. The corresponding salts of the three elements are verysimilar to each other and show a regular variation in properties inpassing from calcium to strontium and from strontium to barium. This isseen in the solubility of the sulphate and hydroxide, and in the ease ofdecomposition of the carbonates, as given in the table. Unlike thealkali metals, their normal carbonates and phosphates are insoluble inwater. CALCIUM ~Occurrence. ~ The compounds of calcium are very abundant in nature, sothat the total amount of calcium in the earth's crust is very large. Agreat many different compounds containing the clement are known, themost important of which are the following: Calcite (marble) CaCO_{3}. Phosphorite Ca_{3}(PO_{4})_{2}. Fluorspar CaF_{2}. Wollastonite CaSiO_{3}. Gypsum CaSO_{4}·2H_{2}O. Anhydrite CaSO_{4}. ~Preparation. ~ Calcium is now prepared by the electrolysis of the meltedchloride, the metal depositing in solid condition on the cathode. It isa gray metal, considerably heavier and harder than sodium. It acts uponwater, forming calcium hydroxide and hydrogen, but the action does notevolve sufficient heat to melt the metal. It promises to become a usefulsubstance, though no commercial applications for it have as yet beenfound. ~Calcium oxide~ (_lime, quicklime_) (CaO). Lime is prepared by stronglyheating calcium carbonate (limestone) in large furnaces called kilns: CaCO_{3} = CaO + CO_{2}. When pure, lime is a white amorphous substance. Heated intensely, as inthe oxyhydrogen flame, it gives a brilliant light called the lime light. Although it is a very difficultly fusible substance, yet in the electricfurnace it can be made to melt and even boil. Water acts upon lime withthe evolution of a great deal of heat, --hence the name quicklime, orlive lime, --the process being called slaking. The equation is CaO + H_{2}O = Ca(OH)_{2}. Lime readily absorbs moisture from the air, and is used to dry moistgases, especially ammonia, which cannot be dried by the usualdesiccating agents. It also absorbs carbon dioxide, forming thecarbonate CaO + CO_{2} = CaCO_{3}. Lime exposed to air is therefore gradually converted into hydroxide andcarbonate, and will no longer slake with water. It is then said to beair-slaked. ~Limekilns. ~ The older kiln, still in common use, consists of a large cylindrical stack in which the limestone is loosely packed. A fire is built at the base of the stack, and when the burning is complete it is allowed to die out and the lime is removed from the kiln. The newer kilns are constructed as shown in Fig. 80. A number of fire boxes are built around the lower part of the kiln, one of which is shown at B. The fire is built on the grate F and the hot products of combustion are drawn up through the stack, decomposing the limestone. The kiln is charged at C, and sometimes fuel is added with the limestone to cause combustion throughout the contents of the kiln. The burned lime is raked out through openings in the bottom of the stack, one of which is shown at _D. _ The advantage of this kind of a kiln over the older form is that the process is continuous, limestone being charged in at the top as fast as the lime is removed at the bottom. [Illustration: Fig. 80] ~Calcium hydroxide ~ (_slaked lime_) (Ca(OH)_{2}). Pure calcium hydroxideis a light white powder. It is sparingly soluble in water, forming asolution called _limewater_, which is often used in medicine as a mildalkali. Chemically, calcium hydroxide is a moderately strong base, though not so strong as sodium hydroxide. Owing to its cheapness it ismuch used in the industries whenever an alkali is desired. A number ofits uses have already been mentioned. It is used in the preparation ofammonia, bleaching powder, and potassium hydroxide. It is also used toremove carbon dioxide and sulphur compounds from coal gas, to remove thehair from hides in the tanneries (this recalls the caustic or corrosiveproperties of sodium hydroxide), and for making mortar. ~Mortar~ is a mixture of calcium hydroxide and sand. When it is exposed tothe air or spread upon porous materials moisture is removed from itpartly by absorption in the porous materials and partly by evaporation, and the mortar becomes firm, or _sets_. At the same time carbon dioxideis slowly absorbed from the air, forming hard calcium carbonate: Ca(OH)_{2} + CO_{2} = CaCO_{3} + H_{2}O. By this combined action the mortar becomes very hard and adheres firmlyto the surface upon which it is spread. The sand serves to give body tothe mortar and makes it porous, so that the change into carbonate cantake place throughout the mass. It also prevents too much shrinkage. ~Cement. ~ When limestone to which clay and sand have been added in certainproportions is burned until it is partly fused (some natural marl isalready of about the right composition), and the clinker so produced isground to powder, the product is called cement. When this material ismoistened it sets to a hard stone-like mass which retains its hardnesseven when exposed to the continued action of water. It can be used forunder-water work, such as bridge piers, where mortar would quicklysoften. Several varieties of cement are made, the best known of which isPortland cement. ~Growing importance of cement. ~ Cement is rapidly coming into use for agreat variety of purposes. It is often used in place of mortar in theconstruction of brick buildings. Mixed with crushed stone and sand itforms concrete which is used in foundation work. It is also used inmaking artificial stone, terra-cotta trimmings for buildings, artificialstone walks and floors, and the like. It is being used more and more formaking many articles which were formerly made of wood or stone, and theentire walls of buildings are sometimes made of cement blocks or ofconcrete. ~Calcium carbonate~ (CaCO_{3}). This substance is found in a great manynatural forms to which various names have been given. They may beclassified under three heads: 1. _Amorphous carbonate. _ This includes those forms which are notmarkedly crystalline. Limestone is the most familiar of these and is agrayish rock usually found in hard stratified masses. Whole mountainranges are sometimes made up of this material. It is always impure, usually containing magnesium carbonate, clay, silica, iron and aluminiumcompounds, and frequently fossil remains. Marl is a mixture of limestoneand clay. Pearls, chalk, coral, and shells are largely calciumcarbonate. 2. _Hexagonal carbonate. _ Calcium carbonate crystallizes in the form ofrhomb-shaped crystals which belong to the hexagonal system. When verypure and transparent the substance is called Iceland spar. Calcite is asimilar form, but somewhat opaque or clouded. Mexican onyx is a massivevariety, streaked or banded with colors due to impurities. Marble whenpure is made up of minute calcite crystals. Stalactites and stalagmitesare icicle-like forms sometimes found in caves. 3. _Rhombic carbonate. _ Calcium carbonate sometimes crystallizes inneedle-shaped crystals belonging to the rhombic system. This is theunstable form and tends to go over into the other variety. Aragonite isthe most familiar example of this form. ~Preparation and uses of calcium carbonate. ~ In the laboratory purecalcium carbonate can be prepared by treating a soluble calcium saltwith a soluble carbonate: Na_{2}CO_{3} + CaCl_{2} = CaCO_{3} + 2NaCl. When prepared in this way it is a soft white powder often calledprecipitated chalk, and is much used as a polishing powder. It isinsoluble in water, but dissolves in water saturated with carbondioxide, owing to the formation of the acid calcium carbonate which isslightly soluble: CaCO_{3} + H_{2}CO_{3} = Ca(HCO_{3})_{2}. The natural varieties of calcium carbonate find many uses, such as inthe preparation of lime and carbon dioxide; in metallurgical operations, especially in the blast furnaces; in the manufacture of soda, glass, andcrayon (which, in addition to chalk, usually contains clay and calciumsulphate); for building stone and ballast for roads. ~Calcium chloride~ (CaCl_{2}). This salt occurs in considerable quantityin sea water. It is obtained as a by-product in many technicalprocesses, as in the Solvay soda process. When crystallized from itssaturated solutions it forms colorless needles of the compositionCaCl_{2}·6H_{2}O. By evaporating a solution to dryness and heating to amoderate temperature calcium chloride is obtained anhydrous as a whiteporous mass. In this condition it absorbs water with great energy and isa valuable drying agent. ~Bleaching powder~ (CaOCl_{2}). When chlorine acts upon a solution ofcalcium hydroxide the reaction is similar to that which occurs betweenchlorine and potassium hydroxide: 2 Ca(OH)_{2} + 4 Cl = CaCl_{2} + Ca(ClO)_{2} + 2 H_{2}O. If, however, chlorine is conducted over calcium hydroxide in the form ofa dry powder, it is absorbed and a substance is formed which appears tohave the composition represented in the formula CaOCl_{2}. Thissubstance is called bleaching powder, or hypochlorite of lime. It isprobably the calcium salt of both hydrochloric and hypochlorous acids, so that its structure is represented by the formula /ClO Ca \Cl. In solution this substance acts exactly like a mixture of calciumchloride (CaCl_{2}) and calcium hypochlorite (Ca(ClO)_{2}), since itdissociates to form the ions Ca^{++}, Cl^{-}, and ClO^{-}. Bleaching powder undergoes a number of reactions which make it animportant substance. 1. When treated with an acid it evolves chlorine: /ClO Ca + H_{2}SO_{4} = CaSO_{4} + HCl + HClO, \Cl HCl + HClO = H_{2}O + 2Cl. This reaction can be employed in the preparation of chlorine, or thenascent chlorine may be used as a bleaching agent. 2. It is slowly decomposed by the carbon dioxide of the air, yieldingcalcium carbonate and chlorine: CaOCl_{2} + CO_{2} = CaCO_{3} + 2Cl. Owing to this slow action the substance is a good disinfectant. 3. When its solution is boiled the substance breaks down into calciumchloride and chlorate: 6CaOCl_{2} = 5CaCl_{2} + Ca(ClO_{3})_{2}. This reaction is used in the preparation of potassium chlorate. ~Calcium fluoride~ (_fluorspar_) (CaF_{2}). Fluorspar has already beenmentioned as the chief natural compound of fluorine. It is found inlarge quantities in a number of localities, and is often crystallized inperfect cubes of a light green or amethyst color. It can be meltedeasily in a furnace, and is sometimes used in the fused condition inmetallurgical operations to protect a metal from the action of the airduring its reduction. It is used as the chief source of fluorinecompounds, especially hydrofluoric acid. ~Calcium sulphate~ (_gypsum_) (CaSO_{4}·2H_{2}O). This abundant substanceoccurs in very perfectly formed crystals or in massive deposits. It isoften found in solution in natural waters and in the sea water. Saltsdeposited from sea water are therefore likely to contain this substance(see Stassfurt salts). It is very sparingly soluble in water, and is thrown down as a finewhite precipitate when any considerable amounts of a calcium salt and asoluble sulphate (or sulphuric acid) are brought together in solution. Its chief use is in the manufacture of plaster of Paris and of hollowtiles for fireproof walls. Such material is called _gypsite_. It is alsoused as a fertilizer. Calcium sulphate, like the carbonate, occurs in many forms in nature. Gypsum is a name given to all common varieties. Granular or massivespecimens are called alabaster, while all those which are wellcrystallized are called selenite. Satin spar is still another varietyoften seen in mineral collections. ~Plaster of Paris. ~ When gypsum is heated to about 115° it loses a portionof its water of crystallization in accordance with the equation 2(CaSO_{4}·2H_{2}O) = 2CaSO_{4}·H_{2}O + 2H_{2}O. The product is a fine white powder called _plaster of Paris_. On beingmoistened it again takes up this water, and in so doing first forms aplastic mass, which soon becomes very firm and hard and regains itscrystalline structure. These properties make it very valuable as amaterial for forming casts and stucco work, for cementing glass tometals, and for other similar purposes. If overheated so that all wateris driven off, the process of taking up water is so slow that thematerial is worthless. Such material is said to be dead burned. Plasterof Paris is very extensively used as the finishing coat for plasteredwalls. ~Hard water. ~ Waters containing compounds of calcium and magnesium insolution are called hard waters because they feel harsh to the touch. The hardness of water may be of two kinds, --(1) temporary hardness and(2) permanent hardness. 1. _Temporary hardness. _ We have seen that when water charged withcarbon dioxide comes in contact with limestone a certain amount of thelatter dissolves, owing to the formation of the soluble acid carbonateof calcium. The hardness of such waters is said to be temporary, sinceit may be removed by boiling. The heat changes the acid carbonate intothe insoluble normal carbonate which then precipitates, rendering thewater soft: Ca(HCO_{3})_{2} = CaCO_{3} + H_{2}O + CO_{2}. Such waters may also be softened by the addition of sufficient lime orcalcium hydroxide to convert the acid carbonate of calcium into thenormal carbonate. The equation representing the reaction is Ca(HCO_{3})_{2} + Ca(OH)_{2} = 2CaCO_{3} + 2H_{2}O. 2. _Permanent hardness. _ The hardness of water may also be due to thepresence of calcium and magnesium sulphates or chlorides. Boiling thewater does not affect these salts; hence such waters are said to havepermanent hardness. They may be softened, however, by the addition ofsodium carbonate, which precipitates the calcium and magnesium asinsoluble carbonates: CaSO_{4} + Na_{2}CO_{3} = CaCO_{3} + Na_{2}SO_{4}. This process is sometimes called "breaking" the water. ~Commercial methods for softening water. ~ The average water of a city supply contains not only the acid carbonates of calcium and magnesium but also the sulphates and chlorides of these metals, together with other salts in smaller quantities. Such waters are softened on a commercial scale by the addition of the proper quantities of calcium hydroxide and sodium carbonate. The calcium hydroxide is added first to precipitate all the acid carbonates. After a short time the sodium carbonate is added to precipitate the other soluble salts of calcium and magnesium, together with any excess of calcium hydroxide which may have been added. The quantity of calcium hydroxide and sodium carbonate required is calculated from a chemical analysis of the water. It will be noticed that the water softened in this way will contain sodium sulphate and chloride, but the presence of these salts is not objectionable. ~Calcium carbide~ (CaC_{2}). This substance is made by heating well-driedcoke and lime in an electrical furnace. The equation is CaO + 3C = CaC_{2} + CO. The pure carbide is a colorless, transparent, crystalline substance. Incontact with water it is decomposed with the evolution of pure acetylenegas, having a pleasant ethereal odor. The commercial article is a dullgray porous substance which contains many impurities. The acetyleneprepared from this substance has a very characteristic odor due toimpurities, the chief of these being phosphine. It is used inconsiderable quantities as a source of acetylene gas for illuminatingpurposes. ~Technical preparation. ~ Fig. 81 represents a recent type of a carbide furnace. The base of the furnace is provided with a large block of carbon A, which serves as one of the electrodes. The other electrodes B, several in number, are arranged horizontally at some distance above this. A mixture of coal and lime is fed into the furnace through the trap top C, and in the lower part of the furnace this mixture becomes intensely heated, forming liquid carbide. This is drawn off through the taphole D. The carbon monoxide formed in the reaction escapes through the pipes E and is led back into the furnace. The pipes F supply air, so that the monoxide burns as it reėnters the furnace and assists in heating the charge. The carbon dioxide so formed, together with the nitrogen entering as air, escape at G. An alternating current is used. [Illustration: Fig. 81] ~Calcium phosphate~ (Ca_{3}(PO_{4})_{2}). This important substanceoccurs abundantly in nature as a constituent of apatite(3Ca_{3}(PO_{4})_{2}·CaF_{2}), in phosphate rock, and as the chiefmineral constituent of bones. Bone ash is therefore nearly pure calciumphosphate. It is a white powder, insoluble in water, although it readilydissolves in acids, being decomposed by them and converted into solubleacid phosphates, as explained in connection with the acids ofphosphorus. STRONTIUM ~Occurrence. ~ Strontium occurs sparingly in nature, usually asstrontianite (SrCO_{3}) and as celestite (SrSO_{4}). Both minerals formbeautiful colorless crystals, though celestite is sometimes colored afaint blue. Only a few of the compounds of strontium have any commercialapplications. ~Strontium hydroxide~ (Sr(OH)_{2}·8H_{2}O). The method of preparation ofstrontium hydroxide is analogous to that of calcium hydroxide. Thesubstance has the property of forming an insoluble compound with sugar, which can easily be separated again into its constituents. It istherefore sometimes used in the sugar refineries to extract sugar fromimpure mother liquors from which the sugar will not crystallize. ~Strontium nitrate~ (Sr(NO_{3})_{2}·4H_{2}O). This salt is prepared bytreating the native carbonate with nitric acid. When ignited withcombustible materials it imparts a brilliant crimson color to the flame, and because of this property it is used in the manufacture of redlights. BARIUM Barium is somewhat more abundant than strontium, occurring in naturelargely as barytes, or heavy spar (BaSO_{4}), and witherite (BaCO_{3}). Like strontium, it closely resembles calcium both in the properties ofthe metal and in the compounds which it forms. ~Oxides of barium. ~ Barium oxide (BaO) can be obtained by strongly heatingthe nitrate: Ba(NO_{3})_{2} = BaO + 2NO_{2} + O. Heated to a low red heat in the air, the oxide combines with oxygen, forming the peroxide (BaO_{2}). If the temperature is raised stillhigher, or the pressure is reduced, oxygen is given off and the oxide isonce more formed. The reaction BaO_{2} BaO + O is reversible and has been used as a means of separating oxygen from theair. Treated with acids, barium peroxide yields hydrogen peroxide: BaO_{2} + 2HCl = BaCl_{2} + H_{2}O_{2}. ~Barium chloride~ (BaCl_{2}·2H_{2}O). Barium chloride is a whitewell-crystallized substance which is easily prepared from the nativecarbonate. It is largely used in the laboratory as a reagent to detectthe presence of sulphuric acid or soluble sulphates. ~Barium sulphate~ _(barytes)_ (BaSO_{4}). Barium sulphate occurs in naturein the form of heavy white crystals. It is precipitated as a crystallinepowder when a barium salt is added to a solution of a sulphate orsulphuric acid: BaCl_{2} + H_{2}SO_{4} = BaSO_{4} + 2HCl. This precipitate is used, as are also the finely ground native sulphateand carbonate, as a pigment in paints. On account of its low cost it issometimes used as an adulterant of white lead, which is also a heavywhite substance. Barium compounds color the flame green, and the nitrate (Ba(NO_{3})_{2})is used in the manufacture of green lights. Soluble barium compounds arepoisonous. RADIUM ~Historical. ~ In 1896 the French scientist Becquerel observed that themineral pitchblende possesses certain remarkable properties. It affectsphotographic plates even in complete darkness, and discharges agold-leaf electroscope when brought close to it. In 1898 Madam Curiemade a careful study of pitchblende to see if these properties belong toit or to some unknown substance contained in it. She succeeded inextracting from it a very small quantity of a substance containing a newelement which she named radium. In 1910 Madam Curie succeeded in obtaining radium itself by theelectrolysis of radium chloride. It is a silver-white metal melting atabout 700°. It blackens in the air, forming a nitride, and decomposeswater. Its atomic weight is about 226. 5. ~Properties. ~ Compounds of radium affect a photographic plate orelectroscope even through layers of paper or sheets of metal. They alsobring about chemical changes in substances placed near them. Investigation of these strange properties has suggested that the radiumatoms are unstable and undergo a decomposition. As a result of thisdecomposition very minute bodies, to which the name corpuscles has beengiven, are projected from the radium atom with exceedingly greatvelocity. It is to these corpuscles that the strange properties ofradium are due. It seems probable that the gas helium is in some wayformed during the decomposition of radium. Two or three other elements, particularly uranium and thorium, have beenfound to possess many of the properties of radium in smaller degree. ~Radium and the atomic theory. ~ If these views in regard to radium shouldprove to be well founded, it will be necessary to modify in somerespects the conception of the atom as developed in a former chapter. The atom would have to be regarded as a compound unit made up of severalparts. In a few cases, as in radium and uranium, it would appear thatthis unit is unstable and undergoes transformation into more stablecombinations. This modification would not, in any essential way, be atvariance with the atomic theory as propounded by Dalton. EXERCISES 1. What properties have the alkaline-earth metals in common with thealkali metals? In what respects do they differ? 2. Write the equation for the reaction between calcium carbide andwater. 3. For what is calcium chlorate used? 4. Could limestone be completely decomposed if heated in a closedvessel? 5. Caves often occur in limestone. Account for their formation. 6. What is the significance of the term fluorspar? (Consult dictionary. ) 7. Could calcium chloride be used in place of barium chloride in testingfor sulphates? 8. What weight of water is necessary to slake the lime obtained from 1ton of pure calcium carbonate? 9. What weight of gypsum is necessary in the preparation of 1 ton ofplaster of Paris? 10. Write equations to represent the reactions involved in thepreparation of strontium hydroxide and strontium nitrate fromstrontianite. 11. Write equations to represent the reactions involved in thepreparation of barium chloride from heavy spar. 12. Could barium hydroxide be used in place of calcium hydroxide intesting for carbon dioxide? CHAPTER XXV THE MAGNESIUM FAMILY =========================================================================== |SYMBOL |ATOMIC |DENSITY |MELTING |BOILING | OXIDE | |WEIGHT | | POINT | POINT |---------------------------------------------------------------------------Magnesium | Mg | 24. 36 | 1. 75 | 750° | 920° | MgOZinc | Zn | 65. 4 | 7. 00 | 420° | 950° | ZnOCadmium | Cd |112. 4 | 8. 67 | 320° | 778° | CdO=========================================================================== ~The family. ~ In the magnesium family are included the four elements:magnesium, zinc, cadmium, and mercury. Between the first three of thesemetals there is a close family resemblance, such as has been tracedbetween the members of the two preceding families. Mercury in somerespects is more similar to copper and will be studied in connectionwith that metal. 1. _Properties. _ When heated to a high temperature in the air each ofthese metals combines with oxygen to form an oxide of the generalformula MO, in which M represents the metal. Magnesium decomposesboiling water slowly, while zinc and cadmium have but little action onit. 2. _Compounds. _ The members of this group are divalent in nearly alltheir compounds, so that the formulas of their salts resemble those ofthe alkaline-earth metals. Like the alkaline-earth metals, theircarbonates and phosphates are insoluble in water. Their sulphates, however, are readily soluble. Unlike both the alkali and alkaline-earthmetals, their hydroxides are nearly insoluble in water. Most of theircompounds dissociate in such a way as to give a simple, colorless, metallic ion. MAGNESIUM ~Occurrence. ~ Magnesium is a very abundant element in nature, ranking alittle below calcium in this respect. Like calcium, it is a constituentof many rocks and also occurs in the form of soluble salts. ~Preparation. ~ The metal magnesium, like most metals whose oxides aredifficult to reduce with carbon, was formerly prepared by heating theanhydrous chloride with sodium: MgCl_{2} + 2Na = 2NaCl + Mg. It is now made by electrolysis, but instead of using as the electrolytethe melted anhydrous chloride, which is difficult to obtain, the naturalmineral carnallite is used. This is melted in an iron pot which alsoserves as the cathode in the electrolysis. A rod of carbon dipping intothe melted salt serves as the anode. The apparatus is very similar tothe one employed in the preparation of sodium. ~Properties. ~ Magnesium is a rather tough silvery-white metal of smalldensity. Air does not act rapidly upon it, but a thin film of oxideforms upon its surface, dimming its bright luster. The common acidsdissolve it with the formation of the corresponding salts. It can beignited readily and in burning liberates much heat and gives a brilliantwhite light. This light is very rich in the rays which affectphotographic plates, and the metal in the form of fine powder isextensively used in the production of flash lights and for white lightsin pyrotechnic displays. ~Magnesium oxide~ (_magnesia_) (MgO). Magnesium oxide, sometimes calledmagnesia or magnesia usta, resembles lime in many respects. It is muchmore easily formed than lime and can be made in the same way, --byigniting the carbonate. It is a white powder, very soft and light, andis unchanged by heat even at very high temperatures. For this reason itis used in the manufacture of crucibles, for lining furnaces, and forother purposes where a refractory substance is needed. It combines withwater to form magnesium hydroxide, but much more slowly and with theproduction of much less heat than in the case of calcium oxide. ~Magnesium hydroxide~ (Mg(OH)_{2}). The hydroxide formed in this way isvery slightly soluble in water, but enough dissolves to give the wateran alkaline reaction. Magnesium hydroxide is therefore a fairly strongbase. It is an amorphous white substance. Neither magnesia nor magnesiumsalts have a very marked effect upon the system; and for this reasonmagnesia is a very suitable antidote for poisoning by strong acids, since any excess introduced into the system will have no injuriouseffect. ~Magnesium cement. ~ A paste of magnesium hydroxide and water slowly absorbs carbon dioxide from the air and becomes very hard. The hardness of the product is increased by the presence of a considerable amount of magnesium chloride in the paste. The hydroxide, with or without the chloride, is used in the preparation of cements for some purposes. ~Magnesium carbonate~ (MgCO_{3}). Magnesium carbonate is a very abundantmineral. It occurs in a number of localities as magnesite, which isusually amorphous, but sometimes forms pure crystals resembling calcite. More commonly it is found associated with calcium carbonate. Themineral dolomite has the composition CaCO_{3}·MgCO_{3}. Limestonecontaining smaller amounts of magnesium carbonate is known as dolomiticlimestone. Dolomite is one of the most common rocks, forming wholemountain masses. It is harder and less readily attacked by acids thanlimestone. It is valuable as a building stone and as ballast forroadbeds and foundations. Like calcium carbonate, magnesium carbonate isinsoluble in water, though easily dissolved by acids. ~Basic carbonate of magnesium. ~ We should expect to find magnesiumcarbonate precipitated when a soluble magnesium salt and a solublecarbonate are brought together: Na_{2}CO_{3} + MgCl_{2} = MgCO_{3} + 2NaCl. Instead of this, some carbon dioxide escapes and the product is found tobe a basic carbonate. The most common basic carbonate of magnesium hasthe formula 4MgCO_{3}·Mg(OH)_{2}, and is sometimes called magnesia alba. This compound is formed by the partial hydrolysis of the normalcarbonate at first precipitated: 5MgCO_{3} + 2H_{2}O = 4MgCO_{3}·Mg(OH)_{2} + H_{2}CO_{3}. ~Magnesium chloride~ (MgCl_{2}·6H_{2}O). Magnesium chloride is found inmany natural waters and in many salt deposits (see Stassfurt salts). Itis obtained as a by-product in the manufacture of potassium chloridefrom carnallite. As there is no very important use for it, largequantities annually go to waste. When heated to drive off the water ofcrystallization the chloride is decomposed as shown in the equation MgCl_{2}·6H_{2}O = MgO + 2HCl + 5H_{2}O. Owing to the abundance of magnesium chloride, this reaction is beingused to some extent in the preparation of both magnesium oxide andhydrochloric acid. ~Boiler scale. ~ When water which contains certain salts in solution is evaporated in steam boilers, a hard insoluble material called _scale_ deposits in the boiler. The formation of this scale may be due to several distinct causes. 1. _To the deposit of calcium sulphate. _ This salt, while sparingly soluble in cold water, is almost completely insoluble in superheated water. Consequently it is precipitated when water containing it is heated in a boiler. 2. _To decomposition of acid carbonates. _ As we have seen, calcium and magnesium acid carbonates are decomposed on heating, forming insoluble normal carbonates: Ca(HCO_{3})_{2} = CaCO_{3} + H_{2}O + CO_{2}. 3. _To hydrolysis of magnesium salts. _ Magnesium chloride, and to some extent magnesium sulphate, undergo hydrolysis when superheated in solution, and the magnesium hydroxide, being sparingly soluble, precipitates: MgCl_{2} + 2H_{2}O Mg(OH)_{2} + 2HCl. This scale adheres tightly to the boiler in compact layers and, being a non-conductor of heat, causes much waste of fuel. It is very difficult to remove, owing to its hardness and resistance to reagents. Thick scale sometimes cracks, and the water coming in contact with the overheated iron occasions an explosion. Moreover, the acids set free in the hydrolysis of the magnesium salts attack the iron tubes and rapidly corrode them. These causes combine to make the formation of scale a matter which occasions much trouble in cases where hard water is used in steam boilers. Water containing such salts should be softened, therefore, before being used in boilers. ~Magnesium sulphate~ (_Epsom salt_) (MgSO_{4}·7H_{2}O). Like the chloride, magnesium sulphate is found rather commonly in springs and in saltdeposits. A very large deposit of the almost pure salt has been found inWyoming. Its name was given to it because of its abundant occurrence inthe waters of the Epsom springs in England. Magnesium sulphate has many uses in the industries. It is used to asmall extent in the preparation of sodium and potassium sulphates, as acoating for cotton cloth, in the dye industry, in tanning, and in themanufacture of paints and laundry soaps. To some extent it is used inmedicine. ~Magnesium silicates. ~ Many silicates containing magnesium are known andsome of them are important substances. Serpentine, asbestos, talc, andmeerschaum are examples of such substances. ZINC ~Occurrence. ~ Zinc never occurs free in nature. Its compounds have beenfound in many different countries, but it is not a constituent of commonrocks and minerals, and its occurrence is rather local and confined todefinite deposits or pockets. It occurs chiefly in the following ores: Sphalerite (zinc blende) ZnS. Zincite ZnO. Smithsonite ZnCO_{3}. Willemite Zn_{2}SiO_{4}. Franklinite ZnO·Fe_{2}O_{3}. One fourth of the world's output of zinc comes from the United States, Missouri being the largest producer. ~Metallurgy. ~ The ores employed in the preparation of zinc are chiefly thesulphide, oxide, and carbonate. They are first roasted in the air, bywhich process they are changed into oxide: ZnCO_{3} = ZnO + CO_{2}, ZnS + 3O = ZnO + SO_{2}. The oxide is then mixed with coal dust, and the mixture is heated inearthenware muffles or retorts, natural gas being used as fuel in manycases. The oxide is reduced by this means to the metallic state, and thezinc, being volatile at the high temperature reached, distills and iscollected in suitable receivers. At first the zinc collects in the formof fine powder, called zinc dust or flowers of zinc, recalling theformation under similar conditions of flowers of sulphur. Later, whenthe whole apparatus has become warm, the zinc condenses to a liquid inthe receiver, from which it is drawn off into molds. Commercial zincoften contains a number of impurities, especially carbon, arsenic, andiron. ~Physical properties. ~ Pure zinc is a rather heavy bluish-white metal witha high luster. It melts at about 420°, and if heated much above thistemperature in the air takes fire and burns with a very bright bluishflame. It boils at about 950° and can therefore be purified bydistillation. Many of the physical properties of zinc are much influenced by thetemperature and previous treatment of the metal. When cast into ingotsfrom the liquid state it becomes at ordinary temperatures quite hard, brittle, and highly crystalline. At 150° it is malleable and can berolled into thin sheets; at higher temperatures it again becomes verybrittle. When once rolled into sheets it retains its softness andmalleability at ordinary temperatures. When melted and poured into waterit forms thin brittle flakes, and in this condition is called granulatedor mossy zinc. ~Chemical properties. ~ Zinc is tarnished superficially by moist air, butbeyond this is not affected by it. It does not decompose even boilingwater. When the metal is quite pure, sulphuric and hydrochloric acidshave scarcely any action upon it; when, however, it contains smallamounts of other metals such as magnesium or arsenic, or when it ismerely in contact with metallic platinum, brisk action takes place andhydrogen is evolved. For this reason, when pure zinc is used in thepreparation of hydrogen a few drops of platinum chloride are often addedto the solution to assist the chemical action. Nitric acid dissolves themetal readily, with the formation of zinc nitrate and various reductionproducts of nitric acid. The strong alkalis act upon zinc and liberatehydrogen: Zn + 2KOH = Zn(OK)_{2} + 2H. The product of this reaction, potassium zincate, is a salt of zinchydroxide, which is thus seen to have acid properties, though it usuallyacts as a base. ~Uses of zinc. ~ The metal has many familiar uses. Rolled into sheets, itis used as a lining for vessels which are to contain water. As a thinfilm upon the surface of iron (galvanized iron) it protects the ironfrom rust. Iron is usually galvanized by dipping it into a bath ofmelted zinc, but electrical methods are also employed. Zinc plates areused in many forms of electrical batteries. In the laboratory zinc isused in the preparation of hydrogen, and in the form of zinc dust as areducing agent. One of the largest uses of zinc is in the manufacture of alloys. Brass, an alloy of zinc and copper, is the most important of these; Germansilver, consisting of copper, zinc, and nickel, has many uses; variousbronzes, coin metals, and bearing metals also contain zinc. Its abilityto alloy with silver finds application in the separation of silver fromlead (see silver). ~Compounds of zinc. ~ In general, the compounds of zinc are similar informula and appearance to those of magnesium, but in other propertiesthey often differ markedly. A number of them have value in commercialways. ~Zinc oxide~ (_zinc white_) (ZnO). Zinc oxide occurs in impure form innature, being colored red by manganese and iron compounds. It can beprepared just like magnesium oxide, but is more often made by burningthe metal. Zinc oxide is a pure white powder which becomes yellow on heating andregains its white color when cold. It is much used as a white pigment inpaints, under the name of zinc white, and has the advantage over whitelead in that it is not changed in color by sulphur compounds, while leadturns black. It is also used in the manufacture of rubber goods. ~Commercial preparation of zinc oxide. ~ Commercially it is often made from franklinite in the following way. The franklinite is mixed with coal and heated to a high temperature in a furnace, by which process the zinc is set free and converted into vapor. As the vapor leaves the furnace through a conduit it meets a current of air and takes fire in it, forming zinc oxide. The oxide passes on and is filtered from the air through canvas bags, which allow the air to pass but retain the oxide. It is thus made by burning the metal, though the metal is not actually isolated in the process. ~Soluble salts. ~ The soluble salts of zinc can be made by dissolving themetal or the oxide in the appropriate acid. They are all somewhatpoisonous. The sulphate and chloride are the most familiar. ~Zinc sulphate~ (_white vitriol_) (ZnSO_{4}·7H_{2}O). This salt is readilycrystallized from strong solutions in transparent colorless crystals. Itis prepared commercially by careful roasting of the sulphide: ZnS + 4O = ZnSO_{4}. ~Zinc chloride~ (ZnCl_{2}·H_{2}O). When a solution of zinc chloride isslowly evaporated a salt of the composition ZnCl_{2}·H_{2}O crystallizesout. If the water is completely expelled by heat and the residuedistilled, the anhydrous chloride is obtained and may be cast intosticks or broken into lumps. In this distillation, just as in heatingmagnesium chloride, some of the chloride is decomposed: ZnCl_{2}·H_{2}O = ZnO + 2HCl. The anhydrous chloride has a great affinity for water, and is used as adehydrating agent. It is also a germicide, and wood which is to beexposed to conditions which favor decay, as, for example, railroad ties, is often soaked in solutions of this salt. ~Insoluble compounds. ~ The insoluble compounds of zinc can be prepared byprecipitation. The most important are the sulphide, carbonate, andhydroxide. ~Zinc sulphide~ (ZnS). This substance occurs as the mineral sphalerite, and is one of the most valued ores of zinc. Very large deposits occur insouthwestern Missouri. The natural mineral is found in large crystals ormasses, resembling resin in color and luster. When prepared byprecipitation the sulphide is white. CADMIUM ~The element. ~ This element occurs in small quantities in some zinc ores. In the course of the metallurgy of zinc the cadmium compounds undergochemical changes quite similar to those of the zinc compounds, and thecadmium distills along with the zinc. Being more volatile, it comes overwith the first of the zinc and is prepared from the first portions ofthe distillate by special methods of purification. The element veryclosely resembles zinc in most respects. Some of its alloys arecharacterized by having low melting points. ~Compounds of cadmium. ~ Among the compounds of cadmium may be mentionedthe chloride (CdCl_{2}·2H_{2}O), the sulphate (3CdSO_{4}·8H_{2}O), andthe nitrate (Cd(NO_{3})_{2}·4H_{2}O). These are white solids soluble inwater. The sulphide (CdS) is a bright yellow substance which isinsoluble in water and in dilute acids. It is valuable as a pigment infine paints. EXERCISES 1. What properties have the metals of the magnesium family in commonwith the alkali metals; with the alkaline-earth metals? 2. Compare the action of the metals of the magnesium group on water withthat of the other metals studied. 3. What metals already studied are prepared by electrolysis? 4. Write the equations representing the reactions between magnesium andhydrochloric acid; between magnesium and dilute sulphuric acid. 5. What property of magnesium was taken advantage of in the isolation ofargon? 6. With phosphoric acid magnesium forms salts similar to those ofcalcium. Write the names and formulas of the corresponding magnesiumsalts. 7. How could you distinguish between magnesium chloride and magnesiumsulphate? between Glauber's salts and Epsom salts? 8. What weight of carnallite is necessary in the preparation of 500 g. Of magnesium? 9. Account for the fact that paints made of zinc oxide are not coloredby hydrosulphuric acid. 10. What hydroxide studied, other than zinc hydroxide, has both acid andbasic properties? 11. Write equations showing how the following compounds of zinc may beobtained from metallic zinc: the oxide, chloride, nitrate, carbonate, sulphate, sulphide, hydroxide. CHAPTER XXVI THE ALUMINIUM FAMILY ~The family. ~ The element aluminium is the most abundant member of thegroup of elements known as the aluminium family; indeed, the othermembers of the family--gallium, indium, and thallium--are of such rareoccurrence that they need not be separately described. The elements ofthe family are ordinarily trivalent, so that the formulas for theircompounds differ from those of the elements so far studied. Theirhydroxides are practically insoluble in water and are very weak bases;indeed, the bases are so weak that their salts are often hydrolyzed intofree base and free acid in solution. The salts formed from these basesusually contain water of crystallization, which cannot be driven offwithout decomposing them more or less. The trivalent metals, which in addition to aluminium include also ironand chromium, are sometimes called the _earth metals_. The name refersto the earthy appearance of the oxides of these metals, and to the factthat many earths, soils, and rocks are composed in part of thesesubstances. ALUMINIUM ~Occurrence. ~ Aluminium never occurs in the free state in nature, owing toits great affinity for oxygen. In combined form, as oxides, silicates, and a few other salts, it is both abundant and widely distributed, beingan essential constituent of all soils and of most rocks exceptinglimestone and sandstone. Cryolite (Na_{3}AlF_{6}), found in Greenland, and bauxite, which is an aluminium hydroxide usually mixed with someiron hydroxide, are important minerals. It is estimated that aluminiumcomposes about 8% of the earth's crust. In the industries the metal iscalled aluminum, but its chemical name is aluminium. [Illustration: Fig. 82] ~Preparation. ~ Aluminium was first prepared by Wöhler, in 1827, by heatinganhydrous aluminium chloride with potassium: AlCl_{3} + 3K = 3KCl + Al. This method was tried after it was found impossible to reduce the oxideof aluminium with carbon. The metal possessed such interestingproperties and promised to be so useful that many efforts were made todevise a cheap way of preparing it. The method which has proved mostsuccessful consists in the electrolysis of the oxide dissolved in meltedcryolite. ~Metallurgy. ~ An iron box A (Fig. 82) about eight feet long and six feet wide is connected with a powerful generator in such a way as to serve as the cathode upon which the aluminium is deposited. Three or four rows of carbon rods B dip into the box and serve as the anodes. The box is partially filled with cryolite and the current is turned on, generating enough heat to melt the cryolite. Aluminium oxide is then added, and under the influence of the electric current it decomposes into aluminium and oxygen. The temperature is maintained above the melting point of aluminium, and the liquid metal, being heavier than cryolite, sinks to the bottom of the vessel, from which it is tapped off from time to time through the tap hole C. The oxygen in part escapes as gas, and in part combines with the carbon of the anode, the combustion being very brilliant. The process is carried on at Niagara Falls. The largest expense in the process, apart from the cost of electrical energy, is the preparation of aluminium oxide free from other oxides, for most of the oxide found in nature is too impure to serve without refining. Bauxite is the principal ore used as a source of the aluminium because it is converted into pure oxide without great difficulty. Since common clay is a silicate of aluminium and is everywhere abundant, it might be expected that this would be utilized in the preparation of aluminium. It is, however, very difficult to extract the aluminium from a silicate, and no practical method has been found which will accomplish this. ~Physical properties. ~ Aluminium is a tin-white metal which melts at 640°and is very light, having a density of 2. 68. It is stiff and strong, andwith frequent annealing can be rolled into thin foil. It is a goodconductor of heat and electricity, though not so good as copper for agiven cross section of wire. ~Chemical properties. ~ Aluminium is not perceptibly acted on by boilingwater, and moist air merely dims its luster. Further action is preventedin each case by the formation of an extremely thin film of oxide uponthe surface of the metal. It combines directly with chlorine, and whenheated in oxygen burns with great energy and the liberation of muchheat. It is therefore a good reducing agent. Hydrochloric acid acts uponit, forming aluminium chloride: nitric acid and dilute sulphuric acidhave almost no action on it, but hot, concentrated sulphuric acid actsupon it in the same way as upon copper: 2Al + 6H_{2}SO_{4} = Al_{2}(SO_{4})_{3} + 6H_{2}O + 3SO_{2}. Alkalis readily attack the metal, liberating hydrogen, as in the case ofzinc: Al + 3KOH = Al(OK)_{3} + 3H. Salt solutions, such as sea water, corrode the metal rapidly. It alloysreadily with other metals. ~Uses of aluminium. ~ These properties suggest many uses for the metal. Itslightness, strength, and permanence make it well adapted for manyconstruction purposes. These same properties have led to its extensiveuse in the manufacture of cooking utensils. The fact that it is easilycorroded by salt solutions is, however, a disadvantage. Owing to itssmall resistance to electrical currents, it is replacing copper to someextent in electrical construction, especially for trolley and powerwires. Some of its alloys have very valuable properties, and aconsiderable part of the aluminium manufactured is used for thispurpose. Aluminium bronze, consisting of about 90% copper and 10%aluminium, has a pure golden color, is strong and malleable, is easilycast, and is permanent in the air. Considerable amounts of aluminiumsteel are also made. ~Goldschmidt reduction process. ~ Aluminium is frequently employed as apowerful reducing agent, many metallic oxides which resist reduction bycarbon being readily reduced by it. The aluminium in the form of a finepowder is mixed with the metallic oxide, together with some substancesuch as fluorspar to act as a flux. The mixture is ignited, and thealuminium unites with the oxygen of the metallic oxide, liberating themetal. This collects in a fused condition under the flux. An enormous quantity of heat is liberated in this reaction, and atemperature as high as 3500° can be reached. The heat of the reaction isturned to practical account in welding car rails, steel castings, and insimilar operations where an intense local heat is required. A mixture ofaluminium with various metallic oxides, ready prepared for suchpurposes, is sold under the name of _thermite_. [Illustration: Fig. 83] ~Preparation of chromium by the Goldschmidt method. ~ A mixture of chromium oxide and aluminium powder is placed in a Hessian crucible (A, Fig. 83), and on top of it is placed a small heap B of a mixture of sodium peroxide and aluminium, into which is stuck a piece of magnesium ribbon C. Powdered fluorspar D is placed around the sodium peroxide, after which the crucible is set on a pan of sand and the magnesium ribbon ignited. When the flame reaches the sodium peroxide mixture combustion of the aluminium begins with almost explosive violence, so that great care must be taken in the experiment. The heat of this combustion starts the reaction in the chromium oxide mixture, and the oxide is reduced to metallic chromium. When the crucible has cooled a button of chromium will be found in the bottom. ~Aluminium oxide~ (Al_{2}O_{3}). This substance occurs in several forms innature. The relatively pure crystals are called corundum, while emery isa variety colored dark gray or black, usually with iron compounds. Intransparent crystals, tinted different colors by traces of impurities, it forms such precious stones as the sapphire, oriental ruby, topaz, andamethyst. All these varieties are very hard, falling little short ofthe diamond in this respect. Chemically pure aluminium oxide can be madeby igniting the hydroxide, when it forms an amorphous white powder: 2Al(OH)_{3} = Al_{2}O_{3} + 3H_{2}O. The natural varieties, corundum and emery, are used for cutting andgrinding purposes; the purest forms, together with the artificiallyprepared oxide, are largely used in the preparation of aluminium. ~Aluminium hydroxide~ (Al(OH)_{3}). The hydroxide occurs in nature as themineral hydrargyllite, and in a partially dehydrated form calledbauxite. It can be prepared by adding ammonium hydroxide to any solublealuminium salt, forming a semi-transparent precipitate which isinsoluble in water but very hard to filter. It dissolves in most acidsto form soluble salts, and in the strong bases to form aluminates, asindicated in the equations Al(OH)_{3} + 3HCl = AlCl_{3} + 3H_{2}O, Al(OH)_{3} + 3NaOH = Al(ONa)_{3} + 3H_{2}O. It may act, therefore, either as a weak base or as a weak acid, itsaction depending upon the character of the substances with which it isin contact. When heated gently the hydroxide loses part of its hydrogenand oxygen according to the equation Al(OH)_{3} = AlO·OH + H_{2}O. This substance, the formula of which is frequently written HAlO_{2}, isa more pronounced acid than is the hydroxide, and its salts arefrequently formed when aluminium compounds are fused with alkalis. Themagnesium salt Mg(AlO_{2})_{2} is called spinel, and many other of itssalts, called aluminates, are found in nature. When heated strongly the hydroxide is changed into oxide, which will notagain take up water on being moistened. ~Mordants and dyeing. ~ Aluminium hydroxide has the peculiar property of combining with many soluble coloring materials and forming insoluble products with them. On this account it is often used as a filter to remove objectionable colors from water. This property also leads to its wide use in the dye industry. Many dyes will not adhere to natural fibers such as cotton and wool, that is, will not "dye fast. " If, however, the cloth to be dyed is soaked in a solution of aluminium compounds and then treated with ammonia, the aluminium salts which have soaked into the fiber will be converted into the hydroxide, which, being insoluble, remains in the body of it. If the fiber is now dipped into a solution of the dye, the aluminium hydroxide combines with the color material and fastens, or "fixes, " it upon the fiber. A substance which serves this purpose is called a _mordant_, and aluminium salts, particularly the acetate, are used in this way. ~Aluminium chloride~ (AlCl_{3}·6 H_{2}O). This substance is prepared bydissolving the hydroxide in hydrochloric acid and evaporating tocrystallization. When heated it is converted into the oxide, resemblingmagnesium in this respect: 2(AlCl_{3}·6 H_{2}O) = Al_{2}O_{3} + 6HCl + 9H_{2}O. The anhydrous chloride, which has some important uses, is made byheating aluminium turnings in a current of chlorine. ~Alums. ~ Aluminium sulphate can be prepared by the action of sulphuricacid upon aluminium hydroxide. It has the property of combining with thesulphates of the alkali metals to form compounds called _alums_. Thus, with potassium sulphate the reaction is expressed by the equation K_{2}SO_{4} + Al_{2}(SO_{4})_{3} + 24H_{2}O = 2(KAl(SO_{4})_{2}·12H_{2}O). Under similar conditions ammonium sulphate yields ammonium alum: (NH_{4})_{2}SO_{4} + Al_{2}(SO_{4})_{3} + 24H_{2}O = 2(NH_{4}Al(SO_{4})_{2}·12H_{2}O). Other trivalent sulphates besides aluminium sulphate can form similarcompounds with the alkali sulphates, and these compounds are also calledalums, though they contain no aluminium. They all crystallize inoctahedra and contain twelve molecules of water of crystallization. Thealums most frequently prepared are the following: Potassium alum KAl(SO_{4})_{2}·12H_{2}O. Ammonium alum NH_{4}Al(SO_{4})_{2}·12H_{2}O. Ammonium iron alum NH_{4}Fe(SO_{4})_{2}·12H_{2}O. Potassium chrome alum KCr(SO_{4})_{2}·12H_{2}O. An alum may therefore be regarded as a compound derived from twomolecules of sulphuric acid, in which one hydrogen atom has beendisplaced by the univalent alkali atom, and the other three hydrogenatoms by an atom of one of the trivalent metals, such as aluminium, iron, or chromium. Very large, well-formed crystals of an alum can be prepared by suspending a small crystal by a thread in a saturated solution of the alum, as shown in Fig. 84. The small crystal slowly grows and assumes a very perfect form. [Illustration: Fig. 84] ~Other salts of aluminium. ~ While aluminium hydroxide forms fairly stablesalts with strong acids, it is such a weak base that its salts with weakacids are readily hydrolyzed. Thus, when an aluminium salt and a solublecarbonate are brought together in solution we should expect to havealuminium carbonate precipitated according to the equation 3Na_{2}CO_{3} + 2AlCl_{3} = Al_{2}(CO_{3})_{3} + 6NaCl. But if it is formed at all, it instantly begins to hydrolyze, theproducts of the hydrolysis being aluminium hydroxide and carbonic acid, Al_{2}(CO_{3})_{3} + 6H_{2}O = 2Al(OH)_{3} + 3H_{2}CO_{3}. Similarly a soluble sulphide, instead of precipitating aluminiumsulphide (Al_{2}S_{3}), precipitates aluminium hydroxide; for hydrogensulphide is such a weak acid that the aluminium sulphide at first formedhydrolyzes at once, forming aluminium hydroxide and hydrogen sulphide: 3Na_{2}S + 2AlCl_{3} + 6H_{2}O = 2Al(OH)_{3} + 6NaCl + 3H_{2}S. ~Alum baking powders. ~ It is because of the hydrolysis of aluminiumcarbonate that alum is used as a constituent of some baking powders. Thealum baking powders consist of a mixture of alum and sodium hydrogencarbonate. When water is added the two compounds react together, formingaluminium carbonate, which hydrolyzes into aluminium hydroxide andcarbonic acid. The carbon dioxide from the latter escapes through thedough and in so doing raises it into a porous condition, which is theend sought in the use of a baking powder. ~Aluminium silicates. ~ One of the most common constituents of rocks isfeldspar (KAlSi_{3}O_{8}), a mixed salt of potassium and aluminium withthe polysilicic acid (H_{4}Si_{3}O_{8}). Under the influence ofmoisture, carbon dioxide, and changes of temperature this substance isconstantly being broken down into soluble potassium compounds andhydrated aluminium silicate. This compound has the formulaAl_{2}Si_{2}O_{7}·2H_{2}O. In relatively pure condition it is calledkaolin; in the impure state, mixed with sand and other substances, itforms common clay. Mica is another very abundant mineral, having varyingcomposition, but being essentially of the formula KAlSiO_{4}. Serpentine, talc, asbestos, and meerschaum are important complexsilicates of aluminium and magnesium, and granite is a mechanicalmixture of quartz, feldspar, and mica. ~Ceramic industries. ~ Many articles of greatest practical importance, ranging from the roughest brick and tile to the finest porcelain and chinaware, are made from some form of kaolin, or clay. No very precise classification of such ware can be made, as the products vary greatly in properties, depending upon the materials used and the treatment during manufacture. Porcelain is made from the purest kaolin, to which must be added some less pure, plastic kaolin, since the pure substance is not sufficiently plastic. There is also added some more fusible substance, such as feldspar, gypsum, or lime, together with some pure quartz. The constituents must be ground very fine, and when thoroughly mixed and moistened must make a plastic mass which can be molded into any desired form. The article molded from such materials is then burned. In this process the article is slowly heated to a point at which it begins to soften and almost fuse, and then it is allowed to cool slowly. At this stage, a very thin vessel will be translucent and have an almost glassy fracture; if, however, it is somewhat thicker, or has not been heated quite so high, it will still be porous, and partly on this account and partly to improve its appearance it is usually glazed. Glazing is accomplished by spreading upon the object a thin layer of a more fusible mixture of the same materials as compose the body of the object itself, and again heating until the glaze melts to a transparent glassy coating upon the surface of the vessel. In some cases fusible mixtures of quite different composition from that used in fashioning the vessel may be used as a glaze. Oxides of lead, zinc, and barium are often used in this way. When less carefully selected materials are used, or quite thick vessels are made, various grades of stoneware are produced. The inferior grades are glazed by throwing a quantity of common salt into the kiln towards the end of the first firing. In the form of vapor the salt attacks the surface of the baked ware and forms an easily fusible sodium silicate upon it, which constitutes a glaze. Vitrified bricks, made from clay or ground shale, are burned until the materials begin to fuse superficially, forming their own glaze. Other forms of brick and tile are not glazed at all, but are left porous. The red color of ordinary brick and earthenware is due to an oxide of iron formed in the burning process. The decorations upon china are sometimes painted upon the baked ware and then glazed over, and sometimes painted upon the glaze and burned in by a third firing. Care must be taken to use such pigments as are not affected by a high heat and do not react chemically with the constituents of the baked ware or the glaze. EXERCISES 1. What metals and compounds studied are prepared by electrolysis? 2. Write the equation for the reaction between aluminium andhydrochloric acid; between aluminium and sulphuric acid (in two steps). 3. What hydroxides other than aluminium hydroxide have both acid andbasic properties? 4. Write equations showing the methods used for preparing aluminiumhydroxide and sulphate. 5. Write the general formula of an alum, representing an atom of analkali metal by X and an atom of a trivalent metal by Y. 6. What is meant by the term polysilicic acid, as used in the discussionof aluminium silicates? 7. Compare the properties of the hydroxides of the different groups ofmetals so far studied. 8. In what respects does aluminium oxide differ from calcium oxide inproperties? 9. Supposing bauxite to be 90% aluminium hydroxide, what weight of it isnecessary for the preparation of 100 kg. Of aluminium? CHAPTER XXVII THE IRON FAMILY =================================================================== | | | | | | | | | APPROXIMATE | | SYMBOL | ATOMIC | DENSITY | MELTING | OXIDES | | WEIGHT | | POINT |________|________|________|_________|_____________|________________ | | | | |Iron | Fe | 55. 9 | 7. 93 | 1800° | FeO, Fe_{2}O_{3}Cobalt | Co | 59. 0 | 8. 55 | 1800° | CoO, Co_{2}O_{3}Nickel | Ni | 58. 7 | 8. 9 | 1600° | NiO, Ni_{2}O_{3}=================================================================== ~The family. ~ The elements iron, cobalt, and nickel form a group in theeighth column of the periodic table. The atomic weights of the three arevery close together, and there is not the same gradual gradation in theproperties of the three elements that is noticed in the families inwhich the atomic weights differ considerably in magnitude. The elementsare very similar in properties, the similarity being so great in thecase of nickel and cobalt that it is difficult to separate them bychemical analysis. The elements occur in nature chiefly as oxides and sulphides, thoughthey have been found in very small quantities in the native state, usually in meteorites. Their sulphides, carbonates, and phosphates areinsoluble in water, the other common salts being soluble. Their saltsare usually highly colored, those of iron being yellow or light green asa rule, those of nickel darker green, while cobalt salts are usuallyrose colored. The metals are obtained by reducing the oxides withcarbon. IRON ~Occurrence. ~ The element iron has long been known, since its ores arevery abundant and it is not difficult to prepare the metal from them infairly pure condition. It occurs in nature in many forms ofcombination, --in large deposits as oxides, sulphides, and carbonates, and in smaller quantities in a great variety of minerals. Indeed, veryfew rocks or soils are free from small amounts of iron, and it isassimilated by plants and animals playing an important part in lifeprocesses. ~Metallurgy. ~ It will be convenient to treat of the metallurgy of ironunder two heads, --Materials Used and Process. ~Materials used. ~ Four distinct materials are used in the metallurgy ofiron: 1. _Iron ore. _ The ores most frequently used in the metallurgy of iron are the following: Hematite Fe_{2}O_{3}. Magnetite Fe_{3}O_{4}. Siderite FeCO_{3}. Limonite 2Fe_{2}O_{2}·3H_{2}O. These ores always contain impurities, such as silica, sulphides, and earthy materials. All ores, with the exception of the oxides, are first roasted to expel any water and carbon dioxide present and to convert any sulphide into oxide. 2. _Carbon. _ Carbon in some form is necessary both as a fuel and as a reducing agent. In former times wood charcoal was used to supply the carbon, but now anthracite coal or coke is almost universally used. 3. _Hot air. _ To maintain the high temperature required for the reduction of iron a very active combustion of fuel is necessary. This is secured by forcing a strong blast of hot air into the lower part of the furnace during the reduction process. 4. _Flux. _ (a) _Purpose of the flux. _ All the materials which enter the furnace must leave it again either in the form of gases or as liquids. The iron is drawn off as the liquid metal after its reduction. To secure the removal of the earthy matter charged into the furnace along with the ore, materials are added to the charge which will, at the high temperature of the furnace, combine with the impurities in the ore, forming a liquid. The material added for this purpose is called the _flux_; the liquid produced from the flux and the ore is called _slag_. (b) _Function of the slag. _ While the main purpose of adding flux to the charge is to remove from the furnace in the form of liquid slag the impurities originally present in the ore, the slag thus produced serves several other functions. It keeps the contents of the furnace in a state of fusion, thus preventing clogging, and makes it possible for the small globules of iron to run together with greater ease into one large liquid mass. (c) _Character of the slag. _ The slag is really a kind of readily fusible glass, being essentially a calcium-aluminium silicate. The ore usually contains silica and some aluminium compounds, so that limestone (which also contains some silica and aluminium) is added to furnish the calcium required for the slag. If the ore and the limestone do not contain a sufficient amount of silica and aluminium for the formation of the slag, these ingredients are added in the form of sand and feldspar. In the formation of slag from these materials the ore is freed from the silica and aluminium which it contained. [Illustration: Fig. 85] ~Process. ~ The reduction of iron is carried out in large towers calledblast furnaces. The blast furnace (Fig. 85) is usually about 80 ft. Highand 20 ft. In internal diameter at its widest part, narrowing somewhatboth toward the top and toward the bottom. The walls are built of steeland lined with fire-brick. The base is provided with a number of pipesT, called tuyers, through which hot air can be forced into thefurnace. The tuyers are supplied from a large pipe S, which circlesthe furnace as a girdle. The base has also an opening M, through whichthe liquid metal can be drawn off from time to time, and a secondopening P, somewhat above the first, through which the excess of slagoverflows. The top is closed by a movable trap C and C, called thecone, and through this the materials to be used are introduced. Thegases produced by the combustion of the fuel and the reduction of theore, together with the nitrogen of the air forced in through the tuyers, escape through pipes D, called downcomer pipes, which leave thefurnace near the top. These gases are very hot and contain combustiblesubstances, principally carbon monoxide; they are therefore utilized asfuel for the engines and also to heat the blast admitted through thetuyers. The lower part of the furnace is often furnished with a waterjacket. This consists of a series of pipes W built into the walls, through which water can be circulated to reduce their temperature. Charges consisting of coke (or anthracite coal), ore, and flux in properproportions are introduced into the furnace at intervals through thetrap top. The coke burns fiercely in the hot-air blast, giving anintense heat and forming carbon monoxide. The ore, working down in thefurnace as the coke burns, becomes very hot, and by the combinedreducing action of the carbon and carbon monoxide is finally reduced tometal and collects as a liquid in the bottom of the furnace, the slagfloating on the molten iron. After a considerable amount of the iron hascollected the slag is drawn off through the opening P. The molten ironis then drawn off into large ladles and taken to the converters for themanufacture of steel, or it is run out into sand molds, forming the barsor ingots called "pigs. " The process is a continuous one, and when oncestarted it is kept in operation for months or even years withoutinterruption. It seems probable that the first product of combustion of the carbon, at the point where the tuyers enter the furnace, is carbon dioxide. This is at once reduced to carbon monoxide by the intensely heated carbon present, so that no carbon dioxide can be found at that point. For practical purposes, therefore, we may consider that carbon monoxide is the first product of combustion. ~Varieties of iron. ~ The iron of commerce is never pure, but containsvarying amounts of other elements, such as carbon, silicon, phosphorus, sulphur, and manganese. These elements may either be alloyed with theiron or may be combined with it in the form of definite chemicalcompounds. In some instances, as in the case of graphite, the mixturemay be merely mechanical. The properties of iron are very much modified by the presence of theseelements and by the form of the combination between them and the iron;the way in which the metal is treated during its preparation has also amarked influence on its properties. Owing to these facts many kinds ofiron are recognized in commerce, the chief varieties being cast iron, wrought iron, and steel. ~Cast iron. ~ The product of the blast furnace, prepared as just described, is called cast iron. It varies considerably in composition, usuallycontaining from 90 to 95% iron, the remainder being largely carbon andsilicon with smaller amounts of phosphorus and sulphur. When the meltedmetal from the blast furnace is allowed to cool rapidly most of thecarbon remains in chemical combination with the iron, and the product iscalled white cast iron. If the cooling goes on slowly, the carbonpartially separates as flakes of graphite which remain scattered throughthe metal. This product is softer and darker in color and is called graycast iron. ~Properties of cast iron. ~ Cast iron is hard, brittle, and rather easilymelted (melting point about 1100°). It cannot be welded or forged intoshape, but is easily cast in sand molds. It is strong and rigid but notelastic. It is used for making castings and in the manufacture of otherkinds of iron. Cast iron, which contains the metal manganese up to theextent of 20%, together with about 3% carbon, is called spiegel iron;when more than this amount of manganese is present the product is calledferromanganese. The ferromanganese may contain as much as 80% manganese. These varieties of cast iron are much used in the manufacture of steel. ~Wrought iron. ~ Wrought iron is made by burning out from cast iron most ofthe carbon, silicon, phosphorus, and sulphur which it contains. Theprocess is called _puddling_, and is carried out in a furnaceconstructed as represented in Fig. 86. The floor of the furnace F issomewhat concave and is made of iron covered with a layer of iron oxide. A long flame produced by burning fuel upon the grate G is directeddownward upon the materials placed upon the floor, and the draught ismaintained by the stack S. A is the ash box and T a trap to catchthe solid particles carried into the stack by the draught. Upon thefloor of the furnace is placed the charge of cast iron, together with asmall amount of material to make a slag. The iron is soon melted by theflame directed upon it, and the sulphur, phosphorus, and silicon areoxidized by the iron oxide, forming oxides which are anhydrides ofacids. These combine with the flux, which is basic in character, or withthe iron oxide, to form a slag. The carbon is also oxidized and escapesas carbon dioxide. As the iron is freed from other elements it becomespasty, owing to the higher melting point of the purer iron, and in thiscondition forms small lumps which are raked together into a larger one. The large lump is then removed from the furnace and rolled or hammeredinto bars, the slag; being squeezed out in this process. The product hasa stranded or fibrous structure. _The product of a puddling furnace iscalled wrought iron. _ [Illustration: Fig. 86] ~Properties of wrought iron. ~ Wrought iron is nearly pure iron, usuallycontaining about 0. 3% of other substances, chiefly carbon. It is tough, malleable, and fibrous in structure. It is easily bent and is notelastic, so it will not sustain pressure as well as cast iron. It can bedrawn out into wire of great tensile strength, and can also be rolledinto thin sheets (sheet iron). It melts at a high temperature (about1600°) and is therefore forged into shape rather than cast. If melted, it would lose its fibrous structure and be changed into a low carbonsteel. ~Steel. ~ Steel, like wrought iron, is made by burning out from cast iron apart of the carbon, silicon, phosphorus, and sulphur which it contains;but the process is carried out in a very different way, and usually, though not always, more carbon is found in steel than in wrought iron. Anumber of processes are in use, but nearly all the steel of commerce ismade by one of the two following methods. [Illustration: Fig. 87] 1. _Bessemer process. _ This process, invented about 1860, is by far themost important. It is carried out in great egg-shaped crucibles calledconverters (Fig. 87), each one of which will hold as much as 15 tons ofsteel. The converter is built of steel and lined with silica. It ismounted on trunnions T, so that it can be tipped over on its side forfilling and emptying. One of the trunnions is hollow and a pipe Pconnects it with an air chamber A, which forms a false bottom to theconverter. The true bottom is perforated, so that air can be forced inby an air blast admitted through the trunnion and the air chamber. White-hot, liquid cast iron from a blast furnace is run into theconverter through its open necklike top O, the converter being tippedover to receive it; the air blast is then turned on and the converterrotated to a nearly vertical position. The elements in the iron arerapidly oxidized, the silicon first and then the carbon. The heatliberated in the oxidation, largely due to the combustion of silicon, keeps the iron in a molten condition. When the carbon is practically allburned out cast iron or spiegel iron, containing a known percentage ofcarbon, is added and allowed to mix thoroughly with the fluid. The steelis then run into molds, and the ingots so formed are hammered or rolledinto rails or other forms. By this process any desired percentage ofcarbon can be added to the steel. Low carbon steel, which does notdiffer much from wrought iron in composition, is now made in this wayand is replacing the more expensive wrought iron for many purposes. ~The basic lining process. ~ When the cast iron contains phosphorus and sulphur in appreciable quantities, the lining of the converter is made of dolomite. The silicon and carbon burn, followed by the phosphorus and sulphur, and the anhydrides of acids so formed combine with the basic oxides of the lining, forming a slag. This is known as the basic lining process. 2. _Open-hearth process. _ In this process a furnace very similar to apuddling furnace is used, but it is lined with silica or dolomiteinstead of iron oxide. A charge consisting in part of old scrap iron ofany kind and in part of cast iron is melted in the furnace by a gasflame. The silicon and carbon are slowly burned away, and when a testshows that the desired percentage of carbon is present the steel is runout of the furnace. _Steel may therefore be defined as the product ofthe Bessemer or open-hearth processes. _ ~Properties of steel. ~ Bessemer and open-hearth steel usually contain onlya few tenths of a per cent of carbon, less than 0. 1% silicon, and a verymuch smaller quantity of phosphorus and sulphur. Any considerable amountof the latter elements makes the steel brittle, the sulphur affecting itwhen hot, and the phosphorus when cold. This kind of steel is used forstructural purposes, for rails, and for nearly all large steel articles. It is hard, malleable, ductile, and melts at a lower temperature thanwrought iron. It can be forged into shape, rolled into sheets, or castin molds. ~Relation of the three varieties of iron. ~ It will be seen that wroughtiron is usually very nearly pure iron, while steel contains anappreciable amount of alloy material, chiefly carbon, and cast ironstill more of the same substances. It is impossible, however, to assigna given sample of iron to one of these three classes on the basis of itschemical composition alone. A low carbon steel, for example, may containless carbon than a given sample of wrought iron. The real distinctionbetween the three is the process by which they are made. The product ofthe blast furnace is cast iron; that of the puddling furnace is wroughtiron; that of the Bessemer and open-hearth methods is steel. ~Tool steel. ~ Steel designed for use in the manufacture of edged tools andsimilar articles should be relatively free from silicon and phosphorus, but should contain from 0. 5 to 1. 5% carbon. The percentage of carbonshould be regulated by the exact use to which the steel is to be put. Steel of this character is usually made in small lots from eitherBessemer or open-hearth steel in the following way. A charge of melted steel is placed in a large crucible and thecalculated quantity of pure carbon is added. The carbon dissolves in thesteel, and when the solution is complete the metal is poured out of thecrucible. This is sometimes called crucible steel. ~Tempering of steel. ~ Steel containing from 0. 5 to 1. 5% carbon ischaracterized by the property of "taking temper. " When the hot steel issuddenly cooled by plunging it into water or oil it becomes very hardand brittle. On carefully reheating this hard form it gradually becomesless brittle and softer, so that by regulating the temperature to whichsteel is reheated in tempering almost any condition of temper demandedfor a given purpose, such as for making springs or cutting tools, can beobtained. ~Steel alloys. ~ It has been found that small quantities of a number ofdifferent elements when alloyed with steel very much improve its qualityfor certain purposes, each element having a somewhat different effect. Among the elements most used in this connection are manganese, silicon, chromium, nickel, tungsten, and molybdenum. The usual method for adding these elements to the steel is to firstprepare a very rich alloy of iron with the element to be added, and thenadd enough of this alloy to a large quantity of the steel to bring it tothe desired composition. A rich alloy of iron with manganese or siliconcan be prepared directly in a blast furnace, and is calledferromanganese or ferrosilicon. Similar alloys of iron with the otherelements mentioned are made in an electric furnace by reducing the mixedoxides with carbon. ~Pure iron. ~ Perfectly pure iron is rarely prepared and is not adapted tocommercial uses. It can be made by reducing pure oxide of iron in acurrent of hydrogen at a high temperature. Prepared in this way itforms a black powder; when melted it forms a tin-white metal which isless fusible and more malleable than wrought iron. It is easily actedupon by moist air. ~Compounds of iron. ~ Iron differs from the metals so far studied in thatit is able to form two series of compounds in which the iron has twodifferent valences. In the one series the iron is divalent and formscompounds which in formulas and many chemical properties are similar tothe corresponding zinc compounds. It can also act as a trivalent metal, and in this condition forms salts similar to those of aluminium. Thosecompounds in which the iron is divalent are known as _ferrous_compounds, while those in which it is trivalent are known as _ferric_. ~Oxides of iron. ~ Iron forms several oxides. Ferrous oxide (FeO) is notfound in nature, but can be prepared artificially in the form of a blackpowder which easily takes up oxygen, forming ferric oxide: 2FeO + O = Fe_{2}O_{3}. Ferric oxide is the most abundant ore of iron and occurs in greatdeposits, especially in the Lake Superior region. It is found in manymineral varieties which vary in density and color, the most abundantbeing hematite, which ranges in color from red to nearly black. Whenprepared by chemical processes it forms a red powder which is used as apaint pigment (Venetian red) and as a polishing powder (rouge). Magnetite has the formula Fe_{3}O_{4} and is a combination of FeO andFe_{2}O_{3}. It is a very valuable ore, but is less abundant thanhematite. It is sometimes called magnetic oxide of iron, or lodestone, since it is a natural magnet. ~Ferrous salts. ~ These salts are obtained by dissolving iron in theappropriate acid, or, when insoluble, by precipitation. They are usuallylight green in color and crystallize well. In chemical reactions theyare quite similar to the salts of magnesium and zinc, but differ fromthem in one important respect, namely, that they are easily changed intocompounds in which the metal is trivalent. Thus ferrous chloride treatedwith chlorine or aqua regia is changed into ferric chloride: FeCl_{2} + Cl = FeCl_{3}. Ferrous hydroxide exposed to moist air is rapidly changed into ferrichydroxide: 2Fe(OH)_{2} + H_{2}O + O = 2Fe(OH)_{3}. ~Ferrous sulphate~ _(copperas, green vitriol)_ (FeSO_{4}·7H_{2}O). Ferroussulphate is the most familiar ferrous compound. It is preparedcommercially as a by-product in the steel-plate mills. Steel plates arecleaned by the action of dilute sulphuric acid upon them, and in theprocess some of the iron dissolves. The liquors are concentrated and thegreen vitriol separates from them. ~Ferrous sulphide~ (FeS). Ferrous sulphide is sometimes found in nature asa golden-yellow crystalline mineral. It is formed as a black precipitatewhen a soluble sulphide and an iron salt are brought together insolution: FeSO_{4} + Na_{2}S = FeS + Na_{2}SO_{4}. It can also be made as a heavy dark-brown solid by fusing together therequisite quantities of sulphur and iron. It is obtained as a by-productin the metallurgy of lead: PbS + Fe = FeS + Pb. It is used in the laboratory in the preparation of hydrosulphuric acid: FeS + 2HCl = FeCl_{2} + H_{2}S. ~Iron disulphide~ _(pyrites)_ (FeS_{2}). This substance bears the samerelation to ferrous sulphide that hydrogen dioxide does to water. Itoccurs abundantly in nature in the form of brass-yellow cubical crystalsand in compact masses. Sometimes the name "fool's gold" is applied to itfrom its superficial resemblance to the precious metal. It is used invery large quantities as a source of sulphur dioxide in the manufactureof sulphuric acid, since it burns readily in the air, forming ferricoxide and sulphur dioxide: 2FeS_{2} + 11O = Fe_{2}O_{3} + 4SO_{2}. ~Ferrous carbonate~ (FeCO_{3}). This compound occurs in nature assiderite, and is a valuable ore. It will dissolve to some extent inwater containing carbon dioxide, just as will calcium carbonate, andwaters containing it are called chalybeate waters. These chalybeatewaters are supposed to possess certain medicinal virtues and form animportant class of mineral waters. ~Ferric salts. ~ Ferric salts are usually obtained by treating an acidifiedsolution of a ferrous salt with an oxidizing agent: 2FeCl_{2} + 2HCl + O = 2FeCl_{3} + H_{2}O, 2FeSO_{4} + H_{2}SO_{4} + O = Fe_{2}(SO_{4})_{3} + H_{2}O. They are usually yellow or violet in color, are quite soluble, and as arule do not crystallize well. Heated with water in the absence of freeacid, they hydrolyze even more readily than the salts of aluminium. Themost familiar ferric salts are the chloride and the sulphate. ~Ferric chloride~ (FeCl_{3}). This salt can be obtained most convenientlyby dissolving iron in hydrochloric acid and then passing chlorine intothe solution: Fe + 2HCl = FeCl_{2} + 2H, FeCl_{2} + Cl = FeCl_{3}. When the pure salt is heated with water it is partly hydrolyzed: FeCl_{3} + 3 H_{2}O Fe(OH)_{3} + 3HCl. This is a reversible reaction, however, and hydrolysis can therefore beprevented by first adding a considerable amount of the soluble productof the reaction, namely, hydrochloric acid. ~Ferric sulphate~ (Fe_{2}(SO_{4})_{3}). This compound can be made bytreating an acid solution of green vitriol with an oxidizing agent. Itis difficult to crystallize and hard to obtain in pure condition. Whenan alkali sulphate in proper quantity is added to ferric sulphate insolution an iron alum is formed, and is easily obtained inlarge crystals. The best known iron alums have the formulasKFe(SO_{4})_{2}·12H_{2}O and NH_{4}Fe(SO_{4})_{2}·12H_{2}O. They arecommonly used when a pure ferric salt is required. ~Ferric hydroxide~ (Fe(OH)_{3}). When solutions of ferric salts aretreated with ammonium hydroxide, ferric hydroxide is formed as arusty-red precipitate, insoluble in water. ~Iron cyanides. ~ A large number of complex cyanides containing iron areknown, the most important being potassium ferrocyanide, or yellowprussiate of potash (K_{4}FeC_{6}N_{6}), and potassium ferricyanide, orred prussiate of potash (K_{3}FeC_{6}N_{6}). These compounds are thepotassium salts of the complex acids of the formulas H_{4}FeC_{6}N_{6}and H_{3}FeC_{6}N_{6}. ~Oxidation of ferrous salts. ~ It has just been seen that when a ferroussalt is treated with an oxidizing agent in the presence of a free acid aferric salt is formed: 2FeSO_{4} + H_{2}SO_{4} + O = Fe_{2}(SO_{4})_{3} + H_{2}O. In this reaction oxygen is used up, and the valence of the iron ischanged from 2 to 3. The same equation may be written 2Fe^{++}, 2SO_{4}^{--} + 2H^{+}, SO_{4}^{--} + O = 2Fe^{+++}, 3SO_{4}^{--} + H_{2}O. Hydrogen ions have been oxidized to water, while the charge of each ironion has been increased from 2 to 3. In a similar way the conversion of ferrous chloride into ferric chloridemay be written Fe^{++}, 2Cl^{-} + Cl = Fe^{+++}, + 3Cl^{-}. Here again the valence of the iron and the charge on the iron ion hasbeen increased from 2 to 3, though no oxygen has entered into thereaction. As a rule, however, changes of this kind are brought about bythe use of an oxidizing agent, and are called oxidations. The term "oxidation" is applied to all reactions in which the valence ofthe metal of a compound is increased, or, in other words, to allreactions in which the charge of a cation is increased. ~Reduction of ferric salts. ~ The changes which take place when a ferricsalt is converted into a ferrous salt are the reverse of the ones justdescribed. This is seen in the equation FeCl_{3} + H = FeCl_{2} + HCl In this reaction the valence of the iron has been changed from 3 to 2. The same equation may be written Fe^{+++}, 3Cl_{-} + H = Fe^{++}, + H^{+} + 3Cl_{-} It will be seen that the charge of the iron ions has been diminishedfrom 3 to 2. Since these changes are the reverse of the oxidationchanges just considered, they are called reduction reactions. The term"reduction" is applied to all processes in which the valence of themetal of a compound is diminished, or, in other words, to all processesin which the charge on the cations is diminished. NICKEL AND COBALT These elements occur sparingly in nature, usually combined with arsenicor with arsenic and sulphur. Both elements have been found in the freestate in meteorites. Like iron they form two series of compounds, butthe salts corresponding to the ferrous salts are the most common, theones corresponding to the ferric salts being difficult to obtain. Thuswe have the chlorides NiCl_{2}·6H_{2}O and CoCl_{2}·6H_{2}O; thesulphates NiSO_{4}·7H_{2}O and CoSO_{4}·7H_{2}O; the nitratesNi(NO_{3})_{2}·6H_{2}O and Co(NO_{3})_{2}·6H_{2}O. Nickel is largely used as an alloy with other metals. Alloyed withcopper it forms coin metal from which five-cent pieces are made, withcopper and zinc it forms German silver, and when added to steel in smallquantities nickel steel is formed which is much superior to common steelfor certain purposes. When deposited by electrolysis upon the surface ofother metals such as iron, it forms a covering which will take a highpolish and protects the metal from rust, nickel not being acted upon bymoist air. Salts of nickel are usually green. Compounds of cobalt fused with glass give it an intensely blue color. Inpowdered form such glass is sometimes used as a pigment called smalt. Cobalt salts, which contain water of crystallization, are usually cherryred in color; when dehydrated they become blue. EXERCISES 1. In the manufacture of cast iron, why is the air heated before beingforced into the furnace? 2. Write the equations showing how each of the following compounds ofiron could be obtained from the metal itself: ferrous chloride, ferroushydroxide, ferrous sulphate, ferrous sulphide, ferrous carbonate, ferricchloride, ferric sulphate, ferric hydroxide. 3. Account for the fact that a solution of sodium carbonate, when addedto a solution of a ferric salt, precipitates an hydroxide and not acarbonate. 4. Calculate the percentage of iron in each of the common iron ores. 5. One ton of steel prepared by the Bessemer process is found byanalysis to contain 0. 2% carbon. What is the minimum weight of carbonwhich must be added in order that the steel may be made to take atemper? CHAPTER XXVIII COPPER, MERCURY, AND SILVER ================================================================== | | | | | | | | | | FORMULAS OF OXIDES | SYMBOL | ATOMIC | DENSITY | MELTING |___________________ | | WEIGHT | | POINT | | | | | | | "ous" | "ic"________|________|________|_________|_________|__________|________ | | | | | |Copper | Cu | 63. 6 | 8. 89 | 1084° | Cu_{2}O | CuOMercury | Hg | 200. 00 | 13. 596 | -39. 5° | Hg_{2}O | HgOSilver | Ag | 107. 93 | 10. 5 | 960° | Ag_{2}O | AgO================================================================== ~The family. ~ By referring to the periodic arrangement of the elements(page 168), it will be seen that mercury is not included in the samefamily with copper and silver. Since the metallurgy of the threeelements is so similar, however, and since they resemble each other soclosely in chemical properties, it is convenient to class them togetherfor study. 1. _Occurrence. _ The three elements occur in nature to some extent inthe free state, but are usually found as sulphides. Their ores are easyto reduce. 2. _Properties. _ They are heavy metals of high luster and are especiallygood conductors of heat and electricity. They are not very activechemically. Neither hydrochloric nor dilute sulphuric acid has anyappreciable action upon them. Concentrated sulphuric acid attacks allthree, forming metallic sulphates and evolving sulphur dioxide, whilenitric acid, both dilute and concentrated, converts them into nitrateswith the evolution of oxides of nitrogen. 3. _Two series of salts. _ Copper and mercury form oxides of the typesM_{2}O and MO, as well as two series of salts. In one series the metalsare univalent and the salts have formulas like those of the sodiumsalts. They are called cuprous and mercurous salts. In the other seriesthe metals are divalent and resemble magnesium salts in formulas. Theseare called cupric and mercuric salts. Silver forms only one series ofsalts, being always a univalent metal. COPPER ~Occurrence. ~ The element copper has been used for various purposes sincethe earliest days of history. It is often found in the metallic state innature, large masses of it occurring pure in the Lake Superior regionand in other places to a smaller extent. The most valuable ores are thefollowing: Cuprite Cu_{2}O. Chalcocite Cu_{2}S. Chalcopyrite CuFeS_{2}. Bornite Cu_{3}FeS_{3}. Malachite CuCO_{3}·Cu(OH)_{2}. Azurite 2CuCO_{3}·Cu(OH)_{2}. ~Metallurgy of copper. ~ Ores containing little or no sulphur are easy toreduce. They are first crushed and the earthy impurities washed away. The concentrated ore is then mixed with carbon and heated in a furnace, metallic copper resulting from the reduction of the copper oxide by thehot carbon. ~Metallurgy of sulphide ores. ~ Much of the copper of commerce is made from chalcopyrite and bornite, and these ores are more difficult to work. They are first roasted in the air, by which treatment much of the sulphur is burned to sulphur dioxide. The roasted ore is then melted in a small blast furnace or in an open one like a puddling furnace. In melting, part of the iron combines with silica to form a slag of iron silicate. The product, called crude matte, contains about 50% copper together with sulphur and iron. Further purification is commonly carried on by a process very similar to the Bessemer process for steel. The converter is lined with silica, and a charge of matte from the melting furnace, together with sand, is introduced, and air is blown into the mass. By this means the sulphur is practically all burned out by the air, and the remaining iron combines with silica and goes off as slag. The copper is poured out of the converter and molded into anode plates for refining. ~Refining of copper. ~ Impure copper is purified by electrolysis. A largeplate of it, serving as an anode, is suspended in a tank facing a thinplate of pure copper, which is the cathode. The tank is filled with asolution of copper sulphate and sulphuric acid to serve as theelectrolyte. A current from a dynamo passes from the anode to thecathode, and the copper, dissolving from the anode, is deposited uponthe cathode in pure form, while the impurities collect on the bottom ofthe tank. Electrolytic copper is one of the purest of commercial metalsand is very nearly pure copper. ~Recovery of gold and silver. ~ Gold and silver are often present in small quantities in copper ores, and in electrolytic refining these metals collect in the muddy deposit on the bottom of the tank. The mud is carefully worked over from time to time and the precious metals extracted from it. A surprising amount of gold and silver is obtained in this way. ~Properties of copper. ~ Copper is a rather heavy metal of density 8. 9, andhas a characteristic reddish color. It is rather soft and is verymalleable, ductile, and flexible, yet tough and strong; it melts at1084°. As a conductor of heat and electrical energy it is second only tosilver. Hydrochloric acid, dilute sulphuric acid, and fused alkalis are almostwithout action upon it; nitric acid and hot, concentrated sulphuricacid, however, readily dissolve it. In moist air it slowly becomescovered with a thin layer of green basic carbonate; heated in the air itis easily oxidized to black copper oxide (CuO). ~Uses. ~ Copper is extensively used for electrical purposes, for roofs andcornices, for sheathing the bottom of ships, and for making alloys. Inthe following table the composition of some of these alloys isindicated: COMPOSITION OF ALLOYS OF COPPER IN PERCENTAGES Aluminium bronze copper (90 to 97%), aluminium (3 to 10%). Brass copper (63 to 73%), zinc (27 to 37%). Bronze copper (70 to 95%), zinc (1 to 25%), tin (1 to 18%). German silver copper (56 to 60%), zinc (20%), nickel (20 to 25%). Gold coin copper (10%), gold (90%). Gun metal copper (90%), tin (10%). Nickel coin copper (75%), nickel (25%) Silver coin copper (10%), silver (90%). ~Electrotyping. ~ Matter is often printed from electrotype plates which are prepared as follows. The matter is set up in type and wax is firmly pressed down upon the face of it until a clear impression is obtained. The impressed side of the wax is coated with graphite and the impression is made the cathode in an electrolytic cell containing a copper salt in solution. When connected with a current the copper is deposited as a thin sheet upon the letters in wax, and when detached is a perfect copy of the type, the under part of the letters being hollow. The sheet is strengthened by pouring on the under surface a suitable amount of molten metal (commercial lead is used). The sheet so strengthened is then used in printing. ~Two series of copper compounds. ~ Copper, like iron, forms two series ofcompounds: in the cuprous compounds it is univalent; in the cupric it isdivalent. The cupric salts are much the more common of the two, sincethe cuprous salts pass readily into cupric by oxidation. ~Cuprous compounds. ~ The most important cuprous compound is the oxide(Cu_{2}O), which occurs in nature as ruby copper or cuprite. It is abright red substance and can easily be prepared by heating copper to ahigh temperature in a limited supply of air. It is used for imparting aruby color to glass. By treating cuprous oxide with different acids a number of cuprous saltscan be made. Many of these are insoluble in water, the chloride (CuCl)being the best known. When suspended in dilute hydrochloric acid it ischanged into cupric chloride, the oxygen taking part in the reactionbeing absorbed from the air: 2CuCl + 2HCl + O = 2CuCl_{2} + H_{2}O. ~Cupric compounds. ~ Cupric salts are easily made by dissolving cupricoxide in acids, or, when insoluble, by precipitation. Most of them areblue or green in color, and the soluble ones crystallize well. Sincethey are so much more familiar than the cuprous salts, they arefrequently called merely copper salts. ~Cupric oxide~ (CuO). This is a black insoluble substance obtained byheating copper in excess of air, or by igniting the hydroxide ornitrate. It is used as an oxidizing agent. ~Cupric hydroxide~ (Cu(OH)_{2}). The hydroxide prepared by treating asolution of a copper salt with sodium hydroxide is a light blueinsoluble substance which easily loses water and changes into the oxide. Heat applied to the liquid containing the hydroxide suspended in itserves to bring about the reaction represented by the equation Cu(OH)_{2} = CuO + H_{2}O. ~Cupric sulphate~ (_blue vitriol_) (CuSO_{4}·5H_{2}O). This substance, called blue vitriol or bluestone, is obtained as a by-product in anumber of processes and is produced in very large quantities. It formslarge blue crystals, which lose water when heated and crumble to a whitepowder. The salt finds many uses, especially in electrotyping and inmaking electrical batteries. ~Cupric sulphide~ (CuS). The insoluble black sulphide (CuS) is easilyprepared by the action of hydrosulphuric acid upon a solution of acopper salt: CuSO_{4} + H_{2}S = CuS + H_{2}SO_{4}. It is insoluble in water and dilute acids. MERCURY ~Occurrence. ~ Mercury occurs in nature chiefly as the sulphide (HgS)called cinnabar, and in globules of metal inclosed in the cinnabar. Themercury mines of Spain have long been famous, California being the nextlargest producer. ~Metallurgy. ~ Mercury is a volatile metal which has but little affinityfor oxygen. Sulphur, on the other hand, readily combines with oxygen. These facts make the metallurgy of mercury very simple. The crushed ore, mixed with a small amount of carbon to reduce any oxide or sulphate thatmight be formed, is roasted in a current of air. The sulphur burns tosulphur dioxide, while the mercury is converted into vapor and iscondensed in a series of condensing vessels. The metal is purified bydistillation. ~Properties. ~ Mercury is a heavy silvery liquid with a density of 13. 596. It boils at 357° and solidifies at -39. 5°. Small quantities of manymetals dissolve in it, forming liquid alloys, while with largerquantities it forms solid alloys. The alloys of mercury are calledamalgams. Toward acids mercury conducts itself very much like copper; it is easilyattacked by nitric and hot, concentrated sulphuric acids, while coldsulphuric and hydrochloric acids have no effect on it. ~Uses. ~ Mercury is extensively used in the construction of scientificinstruments, such as the thermometer and barometer, and as a liquid overwhich to collect gases which are soluble in water. The readiness withwhich it alloys with silver and gold makes it very useful in theextraction of these elements. ~Compounds of mercury. ~ Like copper, mercury forms two series ofcompounds: the mercurous, of which mercurous chloride (HgCl) is anexample; and the mercuric, represented by mercuric chloride (HgCl_{2}). ~Mercuric oxide~ (HgO). Mercuric oxide can be obtained either as abrick-red or as a yellow substance. When mercuric nitrate is heatedcarefully the red modification is formed in accordance with the equation Hg(NO_{3})_{2} = HgO + 2NO_{2} + O. The yellow modification is prepared by adding a solution of a mercuricsalt to a solution of sodium or potassium hydroxide: Hg(NO_{3})_{2} + 2NaOH = 2NaNO_{3} + Hg(OH)_{2}, Hg(OH)_{2} = HgO + H_{2}O. When heated the oxide darkens until it becomes almost black; at a highertemperature it decomposes into mercury and oxygen. It was by thisreaction that oxygen was discovered. ~Mercurous chloride~ (_calomel_) (HgCl). Being insoluble, mercurouschloride is precipitated as a white solid when a soluble chloride isadded to a solution of mercurous nitrate: HgNO_{3} + NaCl = HgCl + NaNO_{3}. Commercially it is manufactured by heating a mixture of mercuricchloride and mercury. When exposed to the light it slowly changes intomercuric chloride and mercury: 2HgCl = HgCl_{2} + Hg. It is therefore protected from the light by the use of colored bottles. It is used in medicine. Most mercurous salts are insoluble in water, the principal soluble onebeing the nitrate, which is made by the action of cold, dilute nitricacid on mercury. ~Mercuric chloride~ (_corrosive sublimate_) (HgCl_{2}). This substance canbe made by dissolving mercuric oxide in hydrochloric acid. On acommercial scale it is made by subliming a mixture of common salt andmercuric sulphate: 2NaCl + HgSO_{4} = HgCl_{2} + Na_{2}SO_{4}. The mercuric chloride, being readily volatile, vaporizes and iscondensed again in cool vessels. Like mercurous chloride it is a whitesolid, but differs from it in that it is soluble in water. It isextremely poisonous and in dilute solutions is used as an antiseptic indressing wounds. ~Mercuric sulphide~ (HgS). As cinnabar this substance forms the chiefnative compound of mercury, occurring in red crystalline masses. Bypassing hydrosulphuric acid into a solution of a mercuric salt it isprecipitated as a black powder, insoluble in water and acids. By othermeans it can be prepared as a brilliant red powder known as vermilion, which is used as a pigment in fine paints. ~The iodides of mercury. ~ If a solution of potassium iodide is added to solutions of a mercurous and a mercuric salt respectively, the corresponding iodides are precipitated. Mercuric iodide is the more important of the two, and as prepared above is a red powder which changes to yellow on heating to 150°. The yellow form on cooling changes back again to the red form, or may be made to do so by rubbing it with a knife blade or some other hard object. SILVER ~Occurrence. ~ Silver is found in small quantities in the uncombined state;usually, however, it occurs in combination with sulphur, either as thesulphide (Ag_{2}S) or as a small constituent of other sulphides, especially those of lead and copper. It is also found alloyed with gold. ~Metallurgy. ~ _Parkes's process. _ Silver is usually smelted in connectionwith lead. The ores are worked over together, as described under lead, and the lead and silver obtained as an alloy, the silver being presentin small quantity. The alloy is melted and metallic zinc is stirred in. Zinc will alloy with silver but not with lead, and it is found that thesilver leaves the lead and, in the form of an alloy with zinc, forms asa crust upon the lead and is skimmed off. This crust, which, of course, contains lead adhering to it, is partially melted and the most of thelead drained off. The zinc is removed by distillation, and the residueis melted on an open hearth in a current of air; by this means the zincand lead remaining with the silver are changed into oxides and thesilver remains behind unaltered. ~Amalgamation process. ~ In some localities the old amalgamation process is used. The silver ore is treated with common salt and ferrous compounds, which process converts the silver first into chloride and then into metallic silver. Mercury is then added and thoroughly mixed with the mass, forming an amalgam with the silver. After some days the earthy materials are washed away and the heavier amalgam is recovered. The mercury is distilled off and the silver left in impure form. ~Refining silver. ~ The silver obtained by either of the above processesmay still contain copper, gold, and iron, and is refined by "parting"with sulphuric acid. The metal is heated with strong sulphuric acidwhich dissolves the silver, copper, and iron present, but not the gold. In the solution of silver sulphate so obtained copper plates aresuspended, upon which the pure silver precipitates, the copper goinginto solution as sulphate, as shown in the equation Ag_{2}SO_{4} + Cu = 2Ag + CuSO_{4}. The solution obtained as a by-product in this process furnishes most ofthe blue vitriol of commerce. Silver is also refined by electrolyticmethods similar to those used in refining copper. ~Properties of silver. ~ Silver is a heavy, rather soft, white metal, veryductile and malleable and capable of taking a high polish. It surpassesall other metals as a conductor of heat and electricity, but is toocostly to find extensive use for such purposes. It melts at a littlelower temperature than copper (961°). It alloys readily with other heavymetals, and when it is to be used for coinage a small amount ofcopper--from 8 to 10%--is nearly always melted with it to give ithardness. It is not acted upon by water or air, but is quickly tarnished when incontact with sulphur compounds, turning quite black in time. Hydrochloric acid and fused alkalis do not act upon it, but nitric acidand hot, concentrated sulphuric acid dissolve it with ease. [Illustration: Fig. 88] ~Electroplating. ~ Since silver is not acted upon by water or air, and has a pleasing appearance, it is used to coat various articles made of cheaper metals. Such articles are said to be silver plated. The process by which this is done is called electroplating. It is carried on as follows: The object to be plated (such as a spoon) is attached to a wire and dipped into a solution of a silver salt. Electrical connection is made in such a way that the article to be plated serves as the cathode, while the anode is made up of one or more plates of silver (Fig. 88, A). When a current is passed through the electrolyte silver dissolves from the anode plate and deposits on the cathode in the form of a closely adhering layer. By making the proper change in the electrolyte and anode plate objects may be plated with gold and other metals. ~Compounds of silver. ~ Silver forms two oxides but only one series ofsalts, namely, the one which corresponds to the mercurous and cuprousseries. ~Silver nitrate~ (_lunar caustic_) (AgNO_{3}). This salt is easilyprepared by dissolving silver in nitric acid and evaporating theresulting solution. It crystallizes in flat plates, and when heatedcarefully can be melted without decomposition. When cast into sticks itis called lunar caustic, for it has a very corrosive action on flesh, and is sometimes used in surgery to burn away abnormal growths. The alchemists designated the metals by the names of the heavenly bodies. The moon (luna) was the symbol for silver; hence the name "lunar caustic. " ~Silver sulphide~ (Ag_{2}S). This occurs in nature and constitutes one ofthe principal ores of silver. It can be obtained in the form of a blacksolid by passing hydrosulphuric acid through a solution of silvernitrate. ~Compounds of silver with the halogens. ~ The chloride, bromide, and iodideof silver are insoluble in water and acids, and are thereforeprecipitated by bringing together a soluble halogen salt with silvernitrate: AgNO_{3} + KCl = AgCl + KNO_{3}. They are remarkable for the fact that they are very sensitive to theaction of light, undergoing a change of color and chemical compositionwhen exposed to sunlight, especially if in contact with organic mattersuch as gelatin. ~Photography. ~ The art of photography is based on the fact that the halogen compounds of silver are affected by the light, particularly in the presence of organic matter. From a chemical standpoint the processes involved may be described under two heads: (1) the preparation of the negative; (2) the preparation of the print. 1. _Preparation of the negative. _ The plate used in the preparation of the negative is made by spreading a thin layer of gelatin, in which silver bromide is suspended (silver iodide is sometimes added also), over a glass plate or celluloid film and allowing it to dry. When the plate so prepared is placed in a camera and the image of some object is focused upon it, the silver salt undergoes a change which is proportional at each point to the intensity of the light falling upon it. In this way an image of the object photographed is produced upon the plate, which is, however, invisible and is therefore called "latent. " It can be made visible by the process of developing. To develop the image the exposed plate is immersed in a solution of some reducing agent called the developer. The developer reduces that portion of the silver salt which has been affected by the light, depositing it in the form of black metallic silver which closely adheres to the plate. The unaffected silver salt, upon which the developer has no action, must now be removed from the plate. This is done by immersing the plate in a solution of sodium thiosulphate (hypo). After the silver salt has been dissolved off, the plate is washed with water and dried. The plate so prepared is called the negative because it is a picture of the object photographed, with the lights exactly reversed. This is called fixing the negative. 2. _Preparation of the print. _ The print is made from paper which is prepared in the same way as the negative plate. The negative is placed upon this paper and exposed to the light in such a way that the light must pass through the negative before striking the paper. If the paper is coated with silver chloride, a visible image is produced, in which case a developer is not needed. The proofs are made in this way. In order to make them permanent the unchanged silver chloride must be dissolved off with sodium thiosulphate. The print is then toned by dipping it into a solution of gold or platinum salts. The silver on the print passes into solution, while the gold or platinum takes its place. These metals give a characteristic color or tone to the print, the gold making it reddish brown, while the platinum gives it a steel-gray tone. If a silver bromide paper is used in making the print, a latent image is produced which must be developed as in the case of the negative itself. The silver bromide is much more sensitive than the chloride, so that the printing can be done in artificial light. Since the darkest places on the negative cut off the most light, it is evident that the lights of the print will be the reverse of those of the negative, and will therefore correspond to those of the object photographed. The print is therefore called the positive. EXERCISES 1. Account for the fact that copper has been used for so long a time. 2. Write equations for the action of concentrated sulphuric and nitricacids upon the metals of this family. 3. How would you account for the fact that normal copper sulphate isslightly acid to litmus? 4. Contrast the action of heat on cupric nitrate and mercuric nitrate. 5. State reasons why mercury is adapted for use in thermometers andbarometers. 6. How could you distinguish between mercurous chloride and mercuricchloride? 7. Write equations for the preparation of mercuric and mercurousiodides. 8. How would you account for the fact that solutions of the differentsalts of a metal usually have the same color? 9. Crude silver usually contains iron and lead. What would become ofthese metals in refining by parting with sulphuric acid? 10. In the amalgamation process for extracting silver, how does ferrouschloride convert silver chloride into silver? Write equation. Why is thesilver sulphide first changed into silver chloride? 11. What impurities would you expect to find in the copper sulphateprepared from the refining of silver? 12. How could you prepare pure silver chloride from a silver coin? 13. Mercuric nitrate and silver nitrate are both white solids soluble inwater. How could you distinguish between them? 14. Account for the fact that sulphur waters turn a silver coin black;also for the fact that a silver spoon is blackened by foods (eggs, forexample) containing sulphur. 15. When a solution of silver nitrate is added to a solution ofpotassium chlorate no precipitate forms. How do you account for the factthat a precipitate of silver chloride is not formed? CHAPTER XXIX TIN AND LEAD ==================================================================== | | | | | | SYMBOL | ATOMIC | DENSITY | MELTING | COMMON OXIDES | | WEIGHT | | POINT |_____|________|________|_________|_________|________________________ | | | | |Tin | Sn | 119. 0 | 7. 35 | 235° | SnO SnO_{2}Lead | Pb | 206. 9 | 11. 38 | 327° | PbO Pb_{3}O_{4} PbO_{2}==================================================================== ~The family. ~ Tin and lead, together with silicon and germanium, form afamily in Group IV of the periodic table. Silicon has been discussedalong with the non-metals, while germanium, on account of its rarity, needs only to be mentioned. The other family of Group IV includes carbon, already described, and anumber of rare elements. TIN ~Occurrence. ~ Tin is found in nature chiefly as the oxide (SnO_{2}), called cassiterite or tinstone. The most famous mines are those ofCornwall in England, and of the Malay Peninsula and East India Islands;in small amounts tinstone is found in many other localities. ~Metallurgy. ~ The metallurgy of tin is very simple. The ore, separated asfar as possible from earthy materials, is mixed with carbon and heatedin a furnace, the reduction taking place readily. The equation is SnO_{2} + C = Sn + CO_{2}. The metal is often purified by carefully heating it until it is partlymelted; the pure tin melts first and can be drained away from theimpurities. ~Properties. ~ Pure tin, called block tin, is a soft white metal with asilver-like appearance and luster; it melts readily (235°) and issomewhat lighter than copper, having a density of 7. 3. It is quitemalleable and can be rolled out into very thin sheets, forming tin foil;most tin foil, however, contains a good deal of lead. Under ordinary conditions it is quite unchanged by air or moisture, butat a high temperature it burns in air, forming the oxide SnO_{2}. Diluteacids have no effect upon it, but concentrated acids attack it readily. Concentrated hydrochloric acid changes it into the chloride Sn + 2HCl = SnCl_{2} + 2H. With sulphuric acid tin sulphate and sulphur dioxide are formed: Sn + 2H_{2}SO_{4} = SnSO_{4} + SO_{2} + 2H_{2}O Concentrated nitric acid oxidizes it, forming a white insoluble compoundof the formula H_{2}SnO_{3}, called metastannic acid: 3Sn + 4HNO_{3} + H_{2}O = 3H_{2}SnO_{3} + 4NO. ~Uses of tin. ~ A great deal of tin is made into tin plate by dipping thinsteel sheets into the melted metal. Owing to the way in which tinresists the action of air and dilute acids, tin plate is used in manyways, such as in roofing, and in the manufacture of tin cans, cookingvessels, and similar articles. Many useful alloys contain tin, some of which have been mentioned inconnection with copper. When tin is alloyed with other metals of lowmelting point, soft, easily melted alloys are formed which are used forfriction bearings in machinery; tin, antimony, lead, and bismuth are thechief constituents of these alloys. Pewter and soft solder are alloys oftin and lead. ~Compounds of tin. ~ Tin forms two series of compounds: the stannous, inwhich the tin is divalent, illustrated in the compounds SnO, SnS, SnCl_{2}; the stannic, in which it is tetravalent as shown in thecompounds SnO_{2}, SnS_{2}. There is also an acid, H_{2}SnO_{3}, calledstannic acid, which forms a series of salts called stannates. While thisacid has the same composition as metastannic acid, the two are quitedifferent in their chemical properties. This difference is probably dueto the different arrangement of the atoms in the molecules of the twosubstances. Only a few compounds of tin need be mentioned. ~Stannic oxide~ (SnO_{2}). Stannic oxide is of interest, since it is thechief compound of tin found in nature. It is sometimes found ingood-sized crystals, but as prepared in the laboratory is a whitepowder. When fused with potassium hydroxide it forms potassium stannate, acting very much like silicon dioxide: SnO_{2} + 2KOH = K_{2}SnO_{3} + H_{2}O. ~Chlorides of tin. ~ Stannous chloride is prepared by dissolving tin inconcentrated hydrochloric acid and evaporating the solution tocrystallization. The crystals which are obtained have the compositionSnCl_{2}·2H_{2}O, and are known as tin crystals. By treating a solutionof stannous chloride with aqua regia, stannic chloride is formed: SnCl_{2} + 2Cl = SnCl_{4}. The salt which crystallizes from such a solution has the compositionSnCl_{4}·5H_{2}O, and is known commercially as oxymuriate of tin. Ifmetallic tin is heated in a current of dry chlorine, the anhydrouschloride (SnCl_{4}) is obtained as a heavy colorless liquid which fumesstrongly on exposure to air. The ease with which stannous chloride takes up chlorine to form stannicchloride makes it a good reducing agent in many reactions, changing thehigher chlorides of metals to lower ones. Thus mercuric chloride ischanged into mercurous chloride: SnCl_{2} + 2HgCl_{2} = SnCl_{4} + 2HgCl. If the stannous chloride is in excess, the reaction may go further, producing metallic mercury: SnCl_{2} + 2HgCl = SnCl_{4} + 2Hg. Ferric chloride is in like manner reduced to ferrous chloride: SnCl_{3} + 2FeCl_{3} = SnCl_{4} + 2FeCl_{2}. The chlorides of tin, as well as the alkali stannates, are much used asmordants in dyeing processes. The hydroxides of tin and free stannicacid, which are easily liberated from these compounds, possess in verymarked degree the power of fixing dyes upon fibers, as explained underaluminium. LEAD ~Occurrence. ~ Lead is found in nature chiefly as the sulphide (PbS), called galena; to a much smaller extent it occurs as carbonate, sulphate, chromate, and in a few other forms. Practically all the leadof commerce is made from galena, two general methods of metallurgy beingin use. ~Metallurgy. ~ 1. The sulphide is melted with scrap iron, when ironsulphide and metallic lead are formed; the liquid lead, being theheavier, sinks to the bottom of the vessel and can be drawn off: PbS + Fe = Pb + FeS. 2. The sulphide is roasted in the air until a part of it has beenchanged into oxide and sulphate. The air is then shut off and theheating continued, the reactions indicated in the following equationstaking place: 2PbO + PbS = 3Pb + SO_{2}, PbSO_{4} + PbS = 2Pb + 2SO_{2}. The lead so prepared usually contains small amounts of silver, arsenic, antimony, copper, and other metals. The silver is removed by Parkes'smethod, as described under silver, and the other metals in various ways. The lead of commerce is one of the purest commercial metals, containingas a rule only a few tenths per cent of impurities. ~Properties. ~ Lead is a heavy metal (den. = 11. 33) which has a brilliantsilvery luster on a freshly cut surface, but which soon tarnishes to adull blue-gray color. It is soft, easily fused (melting at 327°), andquite malleable, but has little toughness or strength. It is not acted upon to any great extent by the oxygen of the air underordinary conditions, but is changed into oxide at a high temperature. With the exception of hydrochloric and sulphuric acids, most acids, evenvery weak ones, act upon it, forming soluble lead salts. Hot, concentrated hydrochloric and sulphuric acids also attack it to a slightextent. ~Uses. ~ Lead is employed in the manufacture of lead pipes and in largestorage batteries. In the form of sheet lead it is used in lining thechambers of sulphuric acid works and in the preparation of paintpigments. Some alloys of lead, such as solder and pewter (lead and tin), shot (lead and arsenic), and soft bearing metals, are widely used. Typemetal consists of lead, antimony, and sometimes tin. Compounds of leadform several important pigments. ~Compounds of lead. ~ In nearly all its compounds lead has a valence of 2, but a few corresponding to stannic compounds have a valence of 4. ~Lead oxides. ~ Lead forms a number of oxides, the most important of whichare litharge, red lead or minium, and lead peroxide. 1. _Litharge_ (PbO). This oxide forms when lead is oxidized at a ratherlow temperature, and is obtained as a by-product in silver refining. Itis a pale yellow powder, and has a number of commercial uses. It iseasily soluble in nitric acid: PbO + 2HNO_{3} = Pb(NO_{3})_{2} + H_{2}O. 2. _Red lead, or minium_ (Pb_{3}O_{4}). Minium is prepared by heatinglead (or litharge) to a high temperature in the air. It is a heavypowder of a beautiful red color, and is much used as a pigment. 3. _Lead peroxide_ (PbO_{2}). This is left as a residue when minium isheated with nitric acid: Pb_{3}O_{4} + 4HNO_{3} = 2Pb(NO_{3})_{2} + PbO_{2} + 2H_{2}O. It is a brown powder which easily gives up a part of its oxygen and, like manganese dioxide and barium dioxide, is a good oxidizing agent. ~Soluble salts of lead. ~ The soluble salts of lead can be made by dissolving(Pb(C_{2}H_{3}O_{2})_{2}·3H_{2}O), litharge in acids. Lead acetatecalled sugar of lead, and lead nitrate (Pb(NO_{3})_{2}) are the mostfamiliar examples. They are while crystalline solids and are poisonousin character. ~Insoluble salts of lead; lead carbonate. ~ While the normal carbonate oflead (PbCO_{3}) is found to some extent, in nature and can be preparedin the laboratory, basic carbonates of varying composition are much moreeasy to obtain. One of the simplest of these has the composition2PbCO_{3}·Pb(OH)_{2}. A mixture of such carbonates is called white lead. This is prepared on a large scale as a paint pigment and as a body forpaints which are to be colored with other substances. ~White lead. ~ White lead is an amorphous white substance which, when mixed with oil, has great covering power, that is, it spreads out in an even waxy film, free from streaks and lumps, and covers the entire surface upon which it is spread. Its disadvantage as a pigment lies in the fact that it gradually blackens when exposed to sulphur compounds, which are often present in the air, forming black lead sulphide (PbS). ~Technical preparation of white lead. ~ Different methods are used in the preparation of white lead, but the old one known as the Dutch process is still the principal one employed. In this process, earthenware pots about ten inches high and of the shape shown in Fig. 89 are used. In the bottom A is placed a 3% solution of acetic acid (vinegar answers the purpose very well). The space above this is filled with thin, perforated, circular pieces of lead, supported by the flange B of the pot. These pots are placed close together on a bed of tan bark on the floor of a room known as the corroding room. They are covered over with boards, upon which tan bark is placed, and another row of pots is placed on this. In this way the room is filled. The white lead is formed by the fumes of the acetic acid, together with the carbon dioxide set free in the fermentation of the tan bark acting on the lead. About three months are required to complete the process. [Illustration 1: Fig. 89] ~Lead sulphide~ (PbS). In nature this compound occurs in highlycrystalline condition, the crystals having much the same luster as purelead. It is readily prepared in the laboratory as a black precipitate, by the action of hydrosulphuric acid upon soluble lead salts: Pb(NO_{3})_{2} + H_{2}S = PbS + 2HNO_{3}. It is insoluble both in water and in dilute acids. ~Other insoluble salts. ~ Lead chromate (PbCrO_{4}) is a yellow substanceproduced by the action of a soluble lead salt upon a soluble chromate, thus: K_{2}CrO_{4} + Pb(NO_{3})_{2} = PbCrO_{4} + 2 KNO_{3}. It is used as a yellow pigment. Lead sulphate (PbSO_{4}) is a whitesubstance sometimes found in nature and easily prepared byprecipitation. Lead chloride (PbCl_{2}) is likewise a white substancenearly insoluble in cold water, but readily soluble in boiling water. ~Thorium and cerium. ~ These elements are found in a few rare minerals, especially in the monazite sand of the Carolinas and Brazil. The oxides of these elements are used in the preparation of the Welsbach mantles for gas lights, because of the intense light given out when a mixture of the oxides is heated. These mantles contain the oxides of cerium and thorium in the ratio of about 1% of the former to 99% of the latter. Compounds of thorium, like those of radium, are found to possess radio-activity, but in a less degree. EXERCISES 1. How could you detect lead if present in tin foil? 2. Stannous chloride reduces gold chloride (AuCl_{3}) to gold. Giveequation. 3. What are the products of hydrolysis when stannic chloride is used asa mordant? 4. How could you detect arsenic, antimony, or copper in lead? 5. Why is lead so extensively used for making water pipes? 6. What sulphates other than lead are insoluble? 7. Could lead nitrate be used in place of barium chloride in testing forsulphates? 8. How much lead peroxide could be obtained from 1 kg. Of minium? 9. The purity of white lead is usually determined by observing thevolume of carbon dioxide given off when it is treated with an acid. Whatacid should be used? On the supposition that it has the formula2PbCO_{3}·Pb(OH)_{2}, how nearly pure was a sample if 1 g. Gave 30 cc. Of carbon dioxide at 20° and 750 mm. ? 10. Silicon belongs in the same family with tin and lead. In whatrespects are these elements similar? 11. What weight of tin could be obtained by the reduction of 1 ton ofcassiterite? 12. What reaction would you expect to take place when lead peroxide istreated with hydrochloric acid? 13. White lead is often adulterated with barytes. Suggest a method fordetecting it, if present, in a given example of white lead. CHAPTER XXX MANGANESE AND CHROMIUM ==================================================================== | | | | | | SYMBOL | ATOMIC | DENSITY | MELTING | FORMULAS OF ACIDS | | WEIGHT | | POINT |__________|________|________|_________|_________|___________________ | | | | |Manganese | Mn | 55. 0 | 8. 01 | 1900° | H_{2}MnO_{4} and | | | | | HMnO_{4}Chromium | Cr | 52. 1 | 7. 3 | 3000° | H_{2}CrO_{4} and | | | | | H_{2}Cr_{2}O_{7}==================================================================== ~General. ~ Manganese and chromium, while belonging to different families, have so many features in common in their chemical conduct that they maybe studied together with advantage. They differ from most of theelements so far studied in that they can act either as acid-forming orbase-forming elements. As base-forming elements each of the metals formstwo series of salts. In the one series, designated by the suffix "ous, "the metal is divalent; in the other series, designated by the suffix"ic, " the metal is trivalent. Only the manganous and the chromic salts, however, are of importance. The acids in which these elements play thepart of a non-metal are unstable, but their salts are usually stable, and some of them are important compounds. MANGANESE ~Occurrence. ~ Manganese is found in nature chiefly as the dioxide MnO_{2}, called pyrolusite. In smaller amounts it occurs as the oxidesMn_{2}O_{3} and Mn_{3}O_{4}, and as the carbonate MnCO_{3}. Some ironores also contain manganese. ~Preparation and properties. ~ The element is difficult to prepare in purecondition and has no commercial applications. It can be prepared, however, by reducing the oxide with aluminium powder or by the use ofthe electric furnace, with carbon as the reducing agent. The metalsomewhat resembles iron in appearance, but is harder, less fusible, andmore readily acted upon by air and moisture. Acids readily dissolve it, forming manganous salts. ~Oxides of manganese. ~ The following oxides of manganese are known: MnO, Mn_{2}O_{3}, Mn_{3}O_{4}, MnO_{2}, and Mn_{2}O_{7}. Only one of these, the dioxide, needs special mention. ~Manganese dioxide~ (_pyrolusite_) (MnO_{2}). This substance is the mostabundant manganese compound found in nature, and is the ore from whichall other compounds of manganese are made. It is a hard, brittle, blacksubstance which is valuable as an oxidizing agent. It will be recalledthat it is used in the preparation of chlorine and oxygen, indecolorizing glass which contains iron, and in the manufacture offerromanganese. ~Compounds containing manganese as a base-forming element. ~ As has beenstated previously, manganese forms two series of salts. The mostimportant of these salts, all of which belong to the manganous series, are the following: Manganous chloride MnCl_{2}·4H_{2}O. Manganous sulphide MnS. Manganous sulphate MnSO_{4}·4H_{2}O. Manganous carbonate MnCO_{3}. Manganous hydroxide Mn(OH)_{2}. The chloride and sulphate may be prepared by heating the dioxide withhydrochloric and sulphuric acids respectively: MnO_{2} + 4HCl = MnCl_{2} + 2H_{2}O + 2Cl, MnO_{2} + H_{2}SO_{4} = MnSO_{4} + H_{2}O + O. The sulphide, carbonate, and hydroxide, being insoluble, may be preparedfrom a solution of the chloride or sulphate by precipitation with theappropriate reagents. Most of the manganous salts are rose colored. Theynot only have formulas similar to the ferrous salts, but resemble themin many of their chemical properties. ~Compounds containing manganese as an acid-forming element. ~ Manganeseforms two unstable acids, namely, manganic acid and permanganic acid. While these acids are of little interest, some of their salts, especially the permanganates, are important compounds. ~Manganic acid and manganates. ~ When manganese dioxide is fused with analkali and an oxidizing agent a green compound is formed. The equation, when caustic potash is used, is as follows: MnO_{2} + 2KOH + O = K_{2}MnO_{4} + H_{2}O. The green compound (K_{2}MnO_{4}) is called potassium manganate, and isa salt of the unstable manganic acid (H_{2}MnO_{4}). The manganates areall very unstable. ~Permanganic acid and the permanganates. ~ When carbon dioxide is passedthrough a solution of a manganate a part of the manganese is changedinto manganese dioxide, while the remainder forms a salt of the unstableacid HMnO_{4}, called permanganic acid. The equation is 3K_{2}MnO_{4} + 2CO_{2} = MnO_{2} + 2KMnO_{4} + 2K_{2}CO_{3}. Potassium permanganate (KMnO_{4}) crystallizes in purple-black needlesand is very soluble in water, forming an intensely purple solution. Allother permanganates, as well as permanganic acid itself, give solutionsof the same color. ~Oxidizing properties of the permanganates. ~ The permanganates areremarkable for their strong oxidizing properties. When used as anoxidizing agent the permanganate is itself reduced, the exact characterof the products formed from it depending upon whether the oxidationtakes place (1) in an alkaline or neutral solution, or (2) in an acidsolution. 1. _Oxidation in alkaline or neutral solution. _ When the solution iseither alkaline or neutral the potassium and the manganese of thepermanganate are both converted into hydroxides, as shown in theequation 2KMnO_{4} + 5H_{2}O = 2Mn(OH)_{4} + 2KOH + 3O. 2. _Oxidation in acid solution. _ When free acid such as sulphuric ispresent, the potassium and the manganese are both changed into salts ofthe acid: 2KMnO_{4} + 3H_{2}SO_{4} = K_{2}SO_{4} + 2MnSO_{4} + 3H_{2}O + 5O. Under ordinary conditions, however, neither one of these reactions takesplace except in the presence of a third substance which is capable ofoxidation. The oxygen is not given off in the free state, as theequations show, but is used up in effecting oxidation. Potassium permanganate is particularly valuable as an oxidizing agentnot only because it acts readily either in acid or in alkaline solution, but also because the reaction takes place so easily that often it is noteven necessary to heat the solution to secure action. The substancefinds many uses in the laboratory, especially in analytical work. It isalso used as an antiseptic as well as a disinfectant. CHROMIUM ~Occurrence. ~ The ore from which all chromium compounds are made ischromite, or chrome iron ore (FeCr_{2}O_{4}). This is found mostabundantly in New Caledonia and Turkey. The element also occurs in smallquantities in many other minerals, especially in crocoisite (PbCrO_{4}), in which mineral it was first discovered. ~Preparation. ~ Chromium, like manganese, is very hard to reduce from itsores, owing to its great affinity for oxygen. It can, however, be madeby the same methods which have proved successful with manganese. Considerable quantities of an alloy of chromium with iron, calledferrochromium, are now produced for the steel industry. ~Properties. ~ Chromium is a very hard metal of about the same density asiron. It is one of the most infusible of the metals, requiring atemperature little short of 3000° for fusion. At ordinary temperaturesair has little action on it; at higher temperatures, however, it burnsbrilliantly. Nitric acid has no action on it, but hydrochloric anddilute sulphuric acids dissolve it, liberating hydrogen. ~Compounds containing chromium as a base-forming element. ~ While chromiumforms two series of salts, chromous salts are difficult to prepare andare of little importance. The most important of the chromic series arethe following: Chromic hydroxide Cr(OH)_{3}. Chromic chloride CrCl_{3}·6H_{2}O. Chromic sulphate Cr_{2}(SO_{4})_{3}. Chrome alums ~Chromic hydroxide~ (Cr(OH)_{3}). This substance, being insoluble, can beobtained by precipitating a solution of the chloride or sulphate with asoluble hydroxide. It is a greenish substance which, like aluminiumhydroxide, dissolves in alkalis, forming soluble salts. ~Dehydration of chromium hydroxide. ~ When heated gently chromic hydroxide loses a part of its oxygen and hydrogen, forming the substance CrO·OH, which, like the corresponding aluminium compound, has more pronounced acid properties than the hydroxide. It forms a series of salts very similar to the spinels; chromite is the ferrous salt of this acid, having the formula Fe(CrO_{2})_{2}. When heated to a higher temperature chromic hydroxide is completely dehydrated, forming the trioxide Cr_{2}O_{3}. This resembles the corresponding oxides of aluminium and iron in many respects. It is a bright green powder, and when ignited strongly becomes almost insoluble in acids, as is also the case with aluminium oxide. ~Chromic sulphate~ (Cr_{2}(SO_{4})_{3}). This compound is a violet-coloredsolid which dissolves in water, forming a solution of the same color. This solution, however, turns green on heating, owing to the formationof basic salts. Chromic sulphate, like ferric and aluminium sulphates, unites with the sulphates of the alkali metals to form alums, of whichthe best known are potassium chrome alum (KCr(SO_{4})_{2}·12H_{2}O) andammonium chrome alum (NH_{4}Cr(SO_{4})_{2}·12H_{2}O). These form beautiful dark purple crystals and have some practical usesin the tanning industry and in photography. A number of the salts ofchromium are also used in the dyeing industry, for they hydrolyze likealuminium salts and the hydroxide forms a good mordant. ~Hydrolysis of chromium salts. ~ When ammonium sulphide is added to a solution of a chromium salt, such as the sulphate, chromium hydroxide precipitates instead of the sulphide. This is due to the fact that chromic sulphide, like aluminium sulphide, hydrolyzes in the presence of water, forming chromic hydroxide and hydrosulphuric acid. Similarly, a soluble carbonate precipitates a basic carbonate of chromium. ~Compounds containing chromium as an acid-forming element. ~ Likemanganese, chromium forms two unstable acids, namely, chromic acid anddichromic acid. Their salts, the chromates and dichromates, areimportant compounds. ~Chromates. ~ When a chromium compound is fused with an alkali and anoxidizing agent a chromate is produced. When potassium hydroxide is usedas the alkali the equation is 2Cr(OH)_{3} + 4KOH + 3O = 2K_{2}CrO_{4} + 5H_{2}O. This reaction recalls the formation of a manganate under similarconditions. ~Properties of chromates. ~ The chromates are salts of the unstable chromicacid (H_{2}CrO_{4}), and as a rule are yellow in color. Lead chromate(PbCrO_{4}) is the well-known pigment chrome yellow. Most of thechromates are insoluble and can therefore be prepared by precipitation. Thus, when a solution of potassium chromate is added to solutions oflead nitrate and barium nitrate respectively, the reactions expressed bythe following equations occur: Pb(NO_{3})_{2} + K_{2}CrO_{4} = PbCrO_{4} + 2KNO_{3}, Ba(NO_{3})_{2} + K_{2}CrO_{4} = BaCrO_{4} + 2KNO_{3}. The chromates of lead and barium separate as yellow precipitates. Thepresence of either of these two metals can be detected by takingadvantage of these reactions. ~Dichromates. ~ When potassium chromate is treated with an acid thepotassium salt of the unstable dichromic acid (H_{2}Cr_{2}O_{7}) isformed: 2K_{2}CrO_{4} + H_{2}SO_{4} = K_{2}Cr_{2}O_{7} + K_{2}SO_{4} + H_{2}O. The relation between the chromates and dichromates is the same as thatbetween the phosphates and the pyrophosphates. Potassium dichromatemight therefore be called potassium pyrochromate. ~Potassium dichromate~ (K_{2}Cr_{2}O_{7}). This is the best knowndichromate, and is the most familiar chromium compound. It forms largecrystals of a brilliant red color, and is rather sparingly soluble inwater. When treated with potassium hydroxide it is converted into thechromate K_{2}Cr_{2}O_{7} + 2KOH = 2K_{2}CrO_{4} + H_{2}O. When added to a solution of lead or barium salt the correspondingchromates (not dichromates) are precipitated. With barium nitrate theequation is 2Ba(NO_{3})_{2} + K_{2}Cr_{2}O_{7} + H_{2}O = 2BaCrO_{4} + 2KNO_{3} + 2HNO_{3}. Potassium dichromate finds use in many industries as an oxidizing agent, especially in the preparation of organic substances, such as the dyealizarin, and in the construction of several varieties of electricbatteries. ~Sodium chromates. ~ The reason why the potassium salt rather than the sodium compound is used is that sodium chromate and dichromate are so soluble that it is hard to prepare them pure. This difficulty is being overcome now, and the sodium compounds are replacing the corresponding potassium salts. This is of advantage, since a sodium salt is cheaper than a potassium salt, so far as raw materials go. ~Oxidizing action of chromates and dichromates. ~ When a dilute solution ofa chromate or dichromate is acidified with an acid, such as sulphuricacid, no reaction apparently takes place. However, if there is present athird substance capable of oxidation, the chromium compound gives up aportion of its oxygen to this substance. Since the chromate changes intoa dichromate in the presence of an acid, it will be sufficient to studythe action of the dichromates alone. The reaction takes place in twosteps. Thus, when a solution of ferrous sulphate is added to a solutionof potassium dichromate acidified with sulphuric acid, the reaction isexpressed by the following equations: (1) K_{2}Cr_{2}O_{7} + 4H_{2}SO_{4} = K_{2}SO_{4} + Cr_{2}(SO_{4})_{3} + 4H_{2}O + 3O, (2) 6FeSO_{4} + 3H_{2}SO_{4} + 3O = 3Fe_{2}(SO_{4})_{3} + 3H_{2}O. The dichromate decomposes in very much the same way as a permanganatedoes, the potassium and chromium being both changed into salts in whichthey play the part of metals, while part of the oxygen of the dichromateis liberated. By combining equations (1) and (2), the following is obtained: K_{2}Cr_{2}O_{7} + 7H_{2}SO_{4} + 6FeSO_{4} = K_{2}SO_{4} + Cr_{2}(SO_{4})_{3} + 3Fe_{2}(SO_{4})_{3} + 7H_{2}0. This reaction is often employed in the estimation of iron in iron ores. ~Potassium chrome alum. ~ It will be noticed that the oxidizing action of potassium dichromate leaves potassium sulphate and chromium sulphate as the products of the reaction. On evaporating the solution these substances crystallize out as potassium chrome alum, which substance is produced as a by-product in the industries using potassium dichromate for oxidizing purposes. ~Chromic anhydride~ (CrO_{3}). When concentrated sulphuric acid is addedto a strong solution of potassium dichromate, and the liquid allowed tostand, deep red needle-shaped crystals appear which have the formulaCrO_{3}. This oxide of chromium is called chromic anhydride, since itcombines readily with water to form chromic acid: CrO_{3} + H_{2}O = H_{2}CrO_{4}. It is therefore analogous to sulphur trioxide which forms sulphuric acidin a similar way: SO_{3} + H_{2}O = H_{2}SO_{4}. Chromic anhydride is a very strong oxidizing agent, giving up oxygen andforming chromic oxide: 2CrO_{3} = Cr_{2}O_{3} + 3O. ~Rare elements of the family. ~ Molybdenum, tungsten, and uranium are three rather rare elements belonging in the same family with chromium, and form many compounds which are similar in formulas to the corresponding compounds of chromium. They can play the part of metals and also form acids resembling chromic acid in formula. Thus we have molybdic acid (H_{2}MoO_{4}), the ammonium salt of which is (NH_{4})_{2}MoO_{4}. This salt has the property of combining with phosphoric acid to form a very complex substance which is insoluble in nitric acid. On this account molybdic acid is often used in the estimation of the phosphoric acid present in a substance. Like chromium, the metals are difficult to prepare in pure condition. Alloys with iron can be prepared by reducing the mixed oxides with carbon in an electric furnace; these alloys are used to some extent in preparing special kinds of steel. EXERCISES 1. How does pyrolusite effect the decolorizing of glass containing iron? 2. Write the equations for the preparation of manganous chloride, carbonate, and hydroxide. 3. Write the equations representing the reactions which take place whenferrous sulphate is oxidized to ferric sulphate by potassiumpermanganate in the presence of sulphuric acid. 4. In the presence of sulphuric acid, oxalic acid is oxidized bypotassium permanganate according to the equation C_{2}H_{2}O_{4} + O = 2CO_{2} + H_{2}O. Write the complete equation. 5. 10 g. Of iron were dissolved in sulphuric acid and oxidized to ferricsulphate by potassium permanganate. What weight of the permanganate wasrequired? 6. What weight of ferrochromium containing 40% chromium must be added toa ton of steel to produce an alloy containing 1% of chromium? 7. Write the equation representing the action of ammonium sulphide uponchromium sulphate. 8. Potassium chromate oxidizes hydrochloric acid, forming chlorine. Write the complete equation. 9. Give the action of sulphuric acid on potassium dichromate (a) inthe presence of a large amount of water; (b) in the presence of asmall amount of water. CHAPTER XXXI GOLD AND THE PLATINUM FAMILY ============================================================================== | | | | | | | | ATOMIC | | HIGHEST | HIGHEST | MELTING | SYMBOL | WEIGHT | DENSITY | OXIDE | CHLORIDE | POINT__________|________|________|_________|_________ |__________|_____________ | | | | | |Ruthenium | Ru | 101. 7 | 12. 26 | RuO_{4} | RuCl_{4} | Electric arcRhodium | Rh | 103. | 12. 1 | RhO_{2} | RhCl_{2} | Electric arcPalladium | Pd | 106. 5 | 11. 8 | PdO_{2} | PdCl_{4} | 1500°Iridium | Ir | 193. | 22. 42 | IrO_{2} | IrCl_{4} | 1950°Osmium | Os | 191. | 22. 47 | OsO_{4} | OsCl_{4} | Electric arcPlatinum | Pt | 194. 8 | 21. 50 | PtO_{2} | PtCl_{4} | 1779°Gold | Au | 197. 2 | 19. 30 | Au_{2}O_{3} | AuCl_{3} | 1064°============================================================================== ~The family. ~ Following iron, nickel, and cobalt in the eighth column ofthe periodic table are two groups of three elements each. The metals ofthe first of these groups--ruthenium, rhodium, and palladium--haveatomic weights near 100 and densities near 12. The metals of the othergroup--iridium, osmium, and platinum--have atomic weights near 200 anddensities near 21. These six rare elements have very similar physicalproperties and resemble each other chemically not only in the type ofcompounds which they form but also in the great variety of them. Theyoccur closely associated in nature, usually as alloys of platinum in theform of irregular metallic grains in sand and gravel. Platinum is by farthe most abundant of the six. Although the periodic classification assigns gold to the silver-coppergroup, its physical as well as many of its chemical properties muchmore closely resemble those of the platinum metals, and it can heconveniently considered along with them. The four elements gold, platinum, osmium, and iridium are the heaviest substances known, beingabout twice as heavy as lead. PLATINUM ~Occurrence. ~ About 90% of the platinum of commerce comes from Russia, small amounts being produced in California, Brazil, and Australia. ~Preparation. ~ Native platinum is usually alloyed with gold and theplatinum metals. To separate the platinum the alloy is dissolved in aquaregia, which converts the platinum into chloroplatinic acid(H_{2}PtCl_{6}). Ammonium chloride is then added, which precipitates theplatinum as insoluble ammonium chloroplatinate: H_{2}PtCl_{6} + 2NH_{4}Cl = (NH_{4})_{2}PtCl_{6} + 2HCl. Some iridium is also precipitated as a similar compound. On ignition thedouble chloride is decomposed, leaving the platinum as a spongy metallicmass, which is melted in an electric furnace and rolled or hammered intothe desired shape. ~Physical properties. ~ Platinum is a grayish-white metal of high luster, and is very malleable and ductile. It melts in the oxyhydrogen blowpipeand in the electric furnace; it is harder than gold and is a goodconductor of electricity. In finely divided form it has the ability toabsorb or occlude gases, especially oxygen and hydrogen. These gases, when occluded, are in a very active condition resembling the nascentstate, and can combine with each other at ordinary temperatures. A jetof hydrogen or coal gas directed upon spongy platinum is at onceignited. ~Platinum as a catalytic agent. ~ Platinum is remarkable for its property of acting as a catalytic agent in a large number of chemical reactions, and mention has been made of this use of the metal in connection with the manufacture of sulphuric acid. When desired for this purpose some porous or fibrous substance, such as asbestos, is soaked in a solution of platinic chloride and then ignited. The platinum compound is decomposed and the platinum deposited in very finely divided form. Asbestos prepared in this way is called platinized asbestos. The catalytic action seems to be in part connected with the property of absorbing gases and rendering them nascent. Some other metals possess this same power, notably palladium, which is remarkable for its ability to absorb hydrogen. ~Chemical properties. ~ Platinum is a very inactive element chemically, andis not attacked by any of the common acids. Aqua regia slowly dissolvesit, forming platinic chloride (PtCl_{4}), which in turn unites with thehydrochloric acid present in the aqua regia, forming the compoundchloroplatinic acid (H_{2}PtCl_{6}). Platinum is attacked by fusedalkalis. It combines at higher temperatures with carbon and phosphorusand alloys with many metals. It is readily attacked by chlorine but notby oxidizing agents. ~Applications. ~ Platinum is very valuable as a material for themanufacture of chemical utensils which are required to stand a hightemperature or the action of strong reagents. Platinum crucibles, dishes, forceps, electrodes, and similar articles are indispensable inthe chemical laboratory. In the industries it is used for such purposesas the manufacture of pans for evaporating sulphuric acid, wires forsealing through incandescent light bulbs, and for making a great varietyof instruments. Unfortunately the supply of the metal is very limited, and the cost is steadily advancing, so that it is now more valuable thangold. ~Compounds. ~ Platinum forms two series of salts of which platinouschloride (PtCl_{2}) and platinic chloride (PtCl_{4}) are examples. Platinates are also known. While a great variety of compounds ofplatinum have been made, the substance is chiefly employed in themetallic state. ~Platinic chloride (PtCl_{4}). ~ Platinic chloride is an orange-colored, soluble compound made by heating chloroplatinic acid in a current ofchlorine. If hydrochloric acid is added to a solution of the substance, the two combine, forming chloroplatinic acid (H_{2}PtCl_{6}): 2HCl + PtCl_{4} = H_{2}PtCl_{6}. The potassium and ammonium salts of this acid are nearly insoluble inwater and alcohol. The acid is therefore used as a reagent toprecipitate potassium in analytical work. With potassium chloride theequation is 2KCl + H_{2}PtCl_{6} = K_{2}PtCl_{6} + 2HCl. ~Other metals of the family. ~ The other members of the family have few applications. Iridium is used in the form of a platinum alloy, since the alloy is much harder than pure platinum and is even less fusible. This alloy is sometimes used to point gold pens. Osmium tetroxide (OsO_{4}) is a very volatile liquid and is used under the name of osmic acid as a stain for sections in microscopy. GOLD ~Occurrence. ~ Gold has been found in many localities, the most famousbeing South Africa, Australia, Russia, and the United States. In thiscountry it is found in Alaska and in nearly half of the states of theunion, notably in California, Colorado, and Nevada. It is usually foundin the native condition, frequently alloyed with silver; in combinationit is sometimes found as telluride (AuTe_{2}), and in a few othercompounds. ~Mining. ~ Native gold occurs in the form of small grains or larger nuggetsin the sands of old rivers, or imbedded in quartz veins in rocks. In thefirst case it is obtained in crude form by placer mining. The sandcontaining the gold is shaken or stirred in troughs of running waterscalled sluices. This sweeps away the sand but allows the heavier gold tosink to the bottom of the sluice. Sometimes the sand containing the goldis washed away from its natural location into the sluices by powerfulstreams of water delivered under pressure from pipes. This is calledhydraulic mining. In vein mining the gold-bearing quartz is mined fromthe veins, stamped into fine powder in stamping mills, and the goldextracted by one of the processes to be described. ~Extraction. ~ 1. _Amalgamation process. _ In the amalgamation process thepowder containing the gold is washed over a series of copper plateswhose surfaces have been amalgamated with mercury. The gold sticks tothe mercury or alloys with it, and after a time the gold and mercury arescraped off and the mixture is distilled. The mercury distills off andthe gold is left in the retort ready for refining. 2. _Chlorination process. _ When gold occurs along with metallicsulphides it is often extracted by chlorination. The ore is firstroasted, and is then moistened and treated with chlorine. This dissolvesthe gold but not the metallic oxides: Au + 3Cl = AuCl_{3}. The gold chloride, being soluble, is extracted from the mixture withwater, and the gold is precipitated from the solution, usually by addingferrous sulphate: AuCl_{3} + 3FeSO_{4} = Au + FeCl_{3} + Fe_{2}(SO_{4})_{3}. 3. _Cyanide process. _ This process depends upon the fact that gold issoluble in a solution of potassium cyanide in the presence of the oxygenof the air. The powder from the stamping mills is treated with a verydilute potassium cyanide solution which extracts the gold: 2Au + 4KCN + H_{2}O + O = 2KOH + 2KAu(CN)_{2}. From this solution the gold can be obtained by electrolysis or byprecipitation with metallic zinc: 2KAu(CN)_{2} + Zn = K_{2}Zn(CN)_{4} + 2Au. ~Refining of gold. ~ Gold is refined by three general methods: 1. _Electrolysis. _ When gold is dissolved in a solution of potassiumcyanide, and the solution electrolyzed, the gold is deposited in verypure condition on the cathode. 2. _Cupellation. _ When the gold is alloyed with easily oxidizablemetals, such as copper or lead, it may be refined by cupellation. Thealloy is fused with an oxidizing flame on a shallow hearth made of boneash, which substance has the property of absorbing metallic oxides butnot the gold. Any silver which may be present remains alloyed with thegold. 3. _Parting with sulphuric acid. _ Gold may be separated from silver, aswell as from many other metals, by heating the alloy with concentratedsulphuric acid. This dissolves the silver, while the gold is notattacked. ~Physical properties. ~ Gold is a very heavy bright yellow metal, exceedingly malleable and ductile, and a good conductor of electricity. It is quite soft and is usually alloyed with copper or silver to give itthe hardness required for most practical uses. The degree of fineness isexpressed in terms of carats, pure gold being twenty-four carats; thegold used for jewelry is usually eighteen carats, eighteen parts beinggold and six parts copper or silver. Gold coinage is 90% gold and 10%copper. ~Chemical properties. ~ Gold is not attacked by any one of the commonacids; aqua regia easily dissolves it, forming gold chloride (AuCl_{3}), which in turn combines with hydrochloric acid to form chlorauric acid(HAuCl_{4}). Fused alkalis also attack it. Most oxidizing agents arewithout action upon it, and in general it is not an active element. ~Compounds. ~ The compounds of gold, though numerous and varied in character, are of comparatively little importance and need not be described in detail. The element forms two series of salts in which it acts as a metal: in the aurous series the gold is univalent, the chloride having the formula AuCl; in the auric series it is trivalent, auric chloride having the formula AuCl_{3}. Gold also acts as an acid-forming element, forming such compounds as potassium aurate (KAuO_{2}). Its compounds are very easily decomposed, however, metallic gold separating from them. EXERCISES 1. From the method of preparation of platinum, what metal is likely tobe alloyed with it? 2. The "platinum chloride" of the laboratory is made by dissolvingplatinum in aqua regia. What is the compound? 3. How would you expect potassium aurate and platinate to be formed?What precautions would this suggest in the use of platinum vessels? 4. Why must gold ores be roasted in the chlorination process? CHAPTER XXXII SOME SIMPLE ORGANIC COMPOUNDS ~Division of chemistry into organic and inorganic. ~ Chemistry is usuallydivided into two great divisions, --organic and inorganic. The originalsignificance of these terms was entirely different from the meaningwhich they have at the present time. 1. _Original significance. _ The division into organic and inorganic wasoriginally made because it was believed that those substances whichconstitute the essential parts of living organisms were built up underthe influence of the life force of the organism. Such substances, therefore, should be regarded as different from those compounds preparedin the laboratory or formed from the inorganic or mineral constituentsof the earth. In accordance with this view organic chemistry includedthose substances formed by living organisms. Inorganic chemistry, on theother hand, included all substances formed from the mineral portions ofthe earth. In 1828 the German chemist Wöhler prepared urea, a typical organiccompound, from inorganic materials. The synthesis of other so-calledorganic compounds followed, and at present it is known that the samechemical laws apply to all substances whether formed in the livingorganism or prepared in the laboratory from inorganic constituents. Theterms "organic" and "inorganic" have therefore lost their originalsignificance. 2. _Present significance. _ The great majority of the compounds found inliving organisms contain carbon, and the term "organic chemistry, " asused at present, includes not only these compounds but all compounds ofcarbon. _Organic chemistry_ has become, therefore, _the chemistry of thecompounds of carbon_, all other substances being treated under the headof inorganic chemistry. This separation of the compounds of carbon intoa group by themselves is made almost necessary by their great number, over one hundred thousand having been recorded. For convenience some ofthe simpler carbon compounds, such as the oxides and the carbonates, areusually discussed in inorganic chemistry. ~The grouping of compounds in classes. ~ The study of organic chemistry ismuch simplified by the fact that the large number of bodies included inthis field may be grouped in classes of similar compounds. It thusbecomes possible to study the properties of each class as a whole, inmuch the same way as we study a group of elements. The most important ofthese classes are the _hydrocarbons_, the _alcohols_, the _aldehydes_, the _acids_, the _ethereal salts_, the _ethers_, the _ketones_, the_organic bases_, and the _carbohydrates_. A few members of each of theseclasses will now be discussed briefly. THE HYDROCARBONS Carbon and hydrogen combine to form a large number of compounds. Thesecompounds are known collectively as the _hydrocarbons_. They may bedivided into a number of groups or series, each being named from itsfirst member. Some of the groups are as follows: METHANE SERIES CH_{4} methane C_{2}H_{6} ethane C_{3}H_{8} propane C_{4}H_{10} butane C_{5}H_{12} pentane C_{6}H_{14} hexane C_{7}H_{16} heptane C_{8}H_{18} octane ETHYLENE SERIES C_{2}H_{4} ethylene C_{3}H_{6} propylene C_{4}H_{8} butylene BENZENE SERIES C_{6}H_{6} benzene C_{7}H_{8} toluene C_{8}H_{10} xylene ACETYLENE SERIES C_{2}H_{2} acetylene C_{3}H_{4} allylene Only the lower members (that is, those which contain a small number ofcarbon atoms) of the above groups are given. The methane series is themost extensive, all of the compounds up to C_{24}H_{50} being known. It will be noticed that the successive members of each of the aboveseries differ by the group of atoms (CH_{2}). Such a series is called an_homologous series_. In general, it may be stated that the members of anhomologous series show a regular gradation in most physical propertiesand are similar in chemical properties. Thus in the methane group thefirst four members are gases at ordinary temperatures; those containingfrom five to sixteen carbon atoms are liquids, the boiling points ofwhich increase with the number of carbon atoms present. Those containingmore than sixteen carbon atoms are solids. ~Sources of the hydrocarbons. ~ There are two chief sources of thehydrocarbons, namely, (1) crude petroleum and (2) coal tar. 1. _Crude petroleum. _ This is a liquid pumped from wells driven into theearth in certain localities. Pennsylvania, Ohio, Kansas, California, andTexas are the chief oil-producing regions in the United States. Thecrude petroleum consists largely of liquid hydrocarbons in which aredissolved both gaseous and solid hydrocarbons. Before being used it mustbe refined. In this process the petroleum is run into large iron stillsand subjected to fractional distillation. The various hydrocarbonsdistill over in the general order of their boiling points. Thedistillates which collect between certain limits of temperature are keptseparate and serve for different uses; they are further purified, generally by washing with sulphuric acid, then with an alkali, andfinally with water. Among the products obtained from crude petroleum inthis way are the naphthas, including benzine and gasoline, kerosene orcoal oil, lubricating oils, vaseline, and paraffin. None of theseproducts are definite chemical compounds, but each consists of a mixtureof hydrocarbons, the boiling points of which lie within certain limits. 2. _Coal tar. _ This product is obtained in the manufacture of coal gas, as already explained. It is a complex mixture and is refined by the samegeneral method used in refining crude petroleum. The principalhydrocarbons obtained from the coal tar are benzene, toluene, naphthalene, and anthracene. In addition to the hydrocarbons, coal tarcontains many other compounds, such as carbolic acid and aniline. ~Properties of the hydrocarbons. ~ The lower members of the first twoseries of hydrocarbons mentioned are all gases; the succeeding membersare liquids. In some series, as the methane series, the higher membersare solids. The preparation and properties of methane and acetylene havebeen discussed in a previous chapter. Ethylene is present in smallquantities in coal gas and may be obtained in the laboratory bytreating alcohol (C_{2}H_{6}O) with sulphuric acid: C_{2}H_{6}O = C_{2}H_{4} + H_{2}O. Benzene, the first member of the benzene series, is a liquid boiling at80°. The hydrocarbons serve as the materials from which a large number ofcompounds can be prepared; indeed, it has been proposed to call organicchemistry _the chemistry of the hydrocarbon derivatives_. ~Substitution products of the hydrocarbons. ~ As a rule, at least a part ofthe hydrogen in any hydrocarbon can be displaced by an equivalent amountof certain elements or groups of elements. Thus the compounds CH_{3}Cl, CH_{2}Cl_{2}, CHCl_{3}, CCl_{4} can be obtained from methane bytreatment with chlorine. Such compounds are called _substitutionproducts_. ~Chloroform~ (CHCl_{3}). This can be made by treating methane withchlorine, as just indicated, although a much easier method consists intreating alcohol or acetone (which see) with bleaching powder. Chloroform is a heavy liquid having a pleasant odor and a sweetishtaste. It is largely used as a solvent and as an anęsthetic in surgery. ~Iodoform~ (CHI_{3}). This is a yellow crystalline solid obtained bytreating alcohol with iodine and an alkali. It has a characteristic odorand is used as an antiseptic. ALCOHOLS When such a compound as CH_{3}Cl is treated with silver hydroxide thereaction expressed by the following equation takes place: CH_{3}Cl + AgOH = CH_{3}OH + AgCl. Similarly C_{2}H_{5}Cl will give C_{2}H_{5}OH and AgCl. The compoundsCH_{3}OH and C_{2}H_{5}OH so obtained belong to the class of substancesknown as _alcohols_. From their formulas it will be seen that they maybe regarded as derived from hydrocarbons by substituting the hydroxylgroup (OH) for hydrogen. Thus the alcohol CH_{3}OH may be regarded asderived from methane (CH_{4}) by substituting the group OH for one atomof hydrogen. A great many alcohols are known, and, like thehydrocarbons, they may be grouped into series. The relation between thefirst three members of the methane series and the corresponding alcoholsis shown in the following table: CH_{4} (methane) CH_{3}OH (methyl alcohol). C_{2}H_{6} (ethane) C_{2}H_{5}OH (ethyl alcohol). C_{3}H_{8} (propane) C_{3}H_{7}OH (propyl alcohol). ~Methyl alcohol~ (_wood alcohol_) (CH_{3}OH). When wood is placed in anair-tight retort and heated, a number of compounds are evolved, the mostimportant of which are the three liquids, methyl alcohol, acetic acid, and acetone. Methyl alcohol is obtained entirely from this source, andon this account is commonly called _wood alcohol_. It is a colorlessliquid which has a density of 0. 79 and boils at 67°. It burns with analmost colorless flame and is sometimes used for heating purposes, inplace of the more expensive ethyl alcohol. It is a good solvent fororganic substances and is used especially as a solvent in themanufacture of varnishes. It is very poisonous. ~Ethyl alcohol~ (_common alcohol_) (C_{2}H_{5}OH). 1. _Preparation. _ Thiscompound may be prepared from glucose (C_{6}H_{12}O_{6}), a sugar easilyobtained from starch. If some baker's yeast is added to a solution ofglucose and the temperature is maintained at about 30°, bubbles of gasare soon evolved, showing that a change is taking place. The yeastcontains a large number of minute organized bodies, which are reallyforms of plant life. The plant grows in the glucose solution, and in sodoing secretes a substance known as _zymase_, which breaks down theglucose in accordance with the following equation: C_{6}H_{12}O_{6} = 2C_{2}H_{5}OH + 2CO_{2}. ~Laboratory preparation of alcohol. ~ The formation of alcohol and carbon dioxide from glucose may be shown as follows: About 100 g. Of glucose are dissolved in a liter of water in flask A (Fig. 90). This flask is connected with the bottle B, which is partially filled with limewater. The tube C contains solid sodium hydroxide. A little baker's yeast is now added to the solution in flask A, and the apparatus is connected, as shown in the figure. If the temperature is maintained at about 30°, the reaction soon begins. The bubbles of gas escape through the limewater in B. A precipitate of calcium carbonate soon forms in the limewater, showing the presence of carbon dioxide. The sodium hydroxide in tube C prevents the carbon dioxide in the air from acting on the limewater. The alcohol remains in the flask A and may be separated by fractional distillation. [Illustration: Fig. 90] 2. _Properties. _ Ethyl alcohol is a colorless liquid with a pleasantodor. It has a density of 0. 78 and boils at 78°. It resembles methylalcohol in its general properties. It is sometimes used as a source ofheat, since its flame is very hot and does not deposit carbon, as theflame from oil does. When taken into the system in small quantities itcauses intoxication; in larger quantities it acts as a poison. Theintoxicating properties of such liquors as beer, wine, and whisky aredue to the alcohol present. Beer contains from 2 to 5% of alcohol, winefrom 5 to 20%, and whisky about 50%. The ordinary alcohol of thedruggist contains 94% of alcohol and 6% of water. When this is boiledwith lime and then distilled nearly all the water is removed, thedistillate being called _absolute alcohol_. ~Commercial preparation of alcohol. ~ Alcohol is prepared commercially from starch obtained from corn or potatoes. The starch is first converted into a sugar known as maltose, by the action of _malt_, a substance prepared by moistening barley with water, allowing it to germinate, and then drying it. There is present in the malt a substance known as diastase, which has the property of changing starch into maltose. This sugar, like glucose, breaks down into alcohol and carbon dioxide in the presence of yeast. The resulting alcohol is separated by fractional distillation. ~Denatured alcohol. ~ The 94% alcohol is prepared at present at a cost of about 35 cents per gallon, which is about half the cost of the preparation of methyl alcohol. The government, however, imposes a tax on all ethyl alcohol which amounts to $2. 08 per gallon on the 94% product. This increases its cost to such an extent that it is not economical to use it for many purposes for which it is adapted, such as a solvent in the preparation of paints and varnishes and as a material for the preparation of many important organic compounds. By an act of Congress in 1906, the tax was removed from _denatured_ alcohol, that is alcohol mixed with some substance which renders it unfit for the purposes of a beverage but will not impair its use for manufacturing purposes. Some of the European countries have similar laws. The substances ordinarily used to denature alcohol are wood alcohol and pyridine, the latter compound having a very offensive odor. ~Fermentation. ~ The reaction which takes place in the preparation of ethyl alcohol belongs to the class of changes known under the general name of fermentation. Thus we say that the yeast causes the glucose to ferment, and the process is known as alcoholic fermentation. There are many kinds of fermentations, and each is thought to be due to the presence of a definite substance known as an _enzyme_, which acts by catalysis. In many cases, as in alcoholic fermentation, the change is brought about by the action of minute forms of life. These probably secrete the enzymes which cause the fermentation to take place. Thus the yeast plant is supposed to bring about alcoholic fermentation by secreting the enzyme known as zymase. ~Glycerin~ (C_{3}H_{5}(OH)_{3}). This compound may be regarded as derivedfrom propane (C_{3}H_{8}) by displacing three atoms of hydrogen by threehydroxyl groups, and must therefore be regarded as an alcohol. It isformed in the manufacture of soaps, as will be explained later. It is anoily, colorless liquid having a sweetish taste. It is used in medicineand in the manufacture of the explosives nitroglycerin and dynamite. ALDEHYDES When alcohols are treated with certain oxidizing agents two hydrogenatoms are removed from each molecule of the alcohol. The resultingcompounds are known as aldehydes. The relation of the aldehydes derivedfrom methyl and ethyl alcohol to the alcohols themselves may be shown asfollows: Alcohols {CH_{3}OH Corresponding aldehydes {CH_{2}O {C_{2}H_{5}OH {C_{2}H_{4}O The first of these (CH_{2}O) is a gas known as formaldehyde. Its aqueoussolution is largely used as an antiseptic and disinfectant under thename of _formalin_. Acetaldehyde (C_{2}H_{4}O) is a liquid boiling at21°. ACIDS Like the other classes of organic compounds, the organic acids may bearranged in homologous series. One of the most important of these seriesis the _fatty-acid series_, the name having been given to it becausethe derivatives of certain of its members are constituents of the fats. Some of the most important members of the series are given in thefollowing table. They are all monobasic, and this fact is expressed inthe formulas by separating the replaceable hydrogen atom from the restof the molecule: H·CHO_{2} formic acid, a liquid boiling at 100°. H·C_{2}H_{3}O acetic acid, a liquid boiling at 118°. H·C_{3}H_{5}O_{2} propionic acid, a liquid boiling at 140°. H·C_{4}H_{7}O_{2} butyric acid, a liquid boiling at 163°. H·C_{16}H_{31}O_{2} palmitic acid, a solid melting at 62°. H·C_{18}H_{35}O_{2} stearic acid, a solid melting at 69°. ~Formic acid~ (H·CHO_{2}). The name "formic" is derived from the Latin_formica_, signifying ant. This name was given to the acid because itwas formerly obtained from a certain kind of ants. It is a colorlessliquid and occurs in many plants such as the stinging nettles. Theinflammation caused by the sting of the bee is due to formic acid. ~Acetic acid~ (H·C_{2}H_{3}O_{2}). Acetic acid is the acid present invinegar, the sour taste being due to it. It can be prepared by either ofthe following methods. 1. _Acetic fermentation. _ This consists in the change of alcohol intoacetic acid through the agency of a minute organism commonly calledmother of vinegar. The change is represented by the following equation: C_{2}H_{5}OH + 2O = HC_{2}H_{3}O_{2} + H_{2}O. The various kinds of vinegars are all made by this process. In themanufacture of cider vinegar the sugar present in the cider firstundergoes alcoholic fermentation; the resulting alcohol then undergoesacetic fermentation. The amount of acetic acid present in vinegarsvaries from 3 to 6%. 2. _From the distillation of wood. _ The liquid obtained by heating woodin the absence of air contains a large amount of acetic acid, and thiscan be separated readily in a pure state. This is the most economicalmethod for the preparation of the concentrated acid. Acetic acid is a colorless liquid and has a strong pungent odor. Many of its salts are well-known compounds. Lead acetate(Pb(C_{2}H_{3}O_{2})_{2}) is the ordinary _sugar of lead_. Sodiumacetate (NaC_{2}H_{3}O_{2}) is a white solid largely used in makingchemical analyses. Copper acetate (Cu(C_{2}H_{3}O_{2})_{2}) is a bluesolid. When copper is acted upon by acetic acid in the presence of air agreen basic acetate of copper is formed. This is commonly known asverdigris. All acetates are soluble in water. ~Butyric acid~ (H·C_{4}H_{7}O_{2}). Derivatives of butyric acid arepresent in butter and impart to it its characteristic flavor. ~Palmitic and stearic acids. ~ Ordinary fats consist principally ofderivatives of palmitic and stearic acids. When the fats are heated withsodium hydroxide the sodium salts of these acids are formed. Ifhydrochloric acid is added to a solution of the sodium salts, the freepalmitic and stearic acids are precipitated. They are white solids, insoluble in water. Stearic acid is often used in making candles. ~Acids belonging to other series. ~ In addition to members of thefatty-acid series, mention may be made of the following well-knownacids. ~Oxalic acid~ (H_{2}C_{2}O_{4}). This is a white solid which occurs innature in many plants, such as the sorrels. Its ammonium salt((NH_{4})_{2}C_{2}O_{4}) is used as a reagent for the detection ofcalcium. When added to a solution of a calcium compound the white, insoluble calcium oxalate (CaC_{2}O_{4}) precipitates. ~Tartaric acid~ (H_{2}·C_{4}H_{4}O_{6}). This compound occurs either in afree state or in the form of its salts in many fruits. The potassiumacid salt (KHC_{4}H_{4}O_{6}) occurs in the juice of grapes. When thejuice ferments in the manufacture of wine, this salt, being insoluble inalcohol, separates out on the sides of the cask and in this form isknown as argol. This is more or less colored by the coloring matter ofthe grape. When purified it forms a white solid and is sold under thename of cream of tartar. The following are also well-known salts oftartaric acid: potassium sodium tartrate (Rochelle salt)(KNaC_{4}H_{4}O_{6}), potassium antimonyl tartrate (tartar emetic)(KSbOC_{4}H_{4}O_{6}). ~Cream of tartar baking powders. ~ The so-called cream of tartar baking powders consist of a mixture of cream of tartar, bicarbonate of soda, and some starch or flour. When water is added to this mixture the cream of tartar slowly acts upon the soda present liberating carbon dioxide in accordance with the following equation: KHC_{4}H_{4}O_{6} + NaHCO_{3} = KNaC_{4}H_{4}O_{6} + H_{2}O + CO_{2}. The carbon dioxide evolved escapes through the dough, thus making it light and porous. ~Citric acid~ (H_{3}·C_{6}H_{5}O_{7}). This acid occurs in many fruits, especially in lemons. It is a white solid, soluble in water, and isoften used as a substitute for lemons in making lemonade. ~Lactic acid~ (H·C_{3}H_{5}O_{3}). This is a liquid which is formed in thesouring of milk. ~Oleic acid~ (H·C_{18}H_{33}O_{2}). The derivatives of this acidconstitute the principal part of many oils and liquid fats. The aciditself is an oily liquid. ETHEREAL SALTS When acids are brought in contact with alcohols under certain conditionsa reaction takes place similar to that which takes place between acidsand bases. The following equations will serve as illustrations: KOH + HNO_{3} = KNO_{3} + H_{2}O, CH_{3}OH + HNO_{3} = CH_{3}NO_{3} + H_{2}O. The resulting compounds of which methyl nitrate (CH_{3}NO_{3}) may betaken as the type belong to the class known as _ethereal salts_, thename having been given them because some of them possess pleasantethereal odors. It will be seen that the ethereal salts differ fromordinary salts in that they contain a hydrocarbon radical, such asCH_{3}, C_{2}H_{5}, C_{3}H_{5}, in place of a metal. ~The nitrates of glycerin~ (_nitroglycerin_). Nitric acid reacts withglycerin in the same way that it reacts with a base containing threehydroxyl groups such as Fe(OH)_{3}: Fe(OH)_{3} + 3HNO_{3} = Fe(NO_{3})_{3} + 3H_{2}O, C_{3}H_{5}(OH)_{3} + 3HNO_{3} = C_{3}H_{5}(NO_{3})_{3} + 3H_{2}O. The resulting nitrate (C_{3}H_{5}(NO_{3})_{3}) is the main constituentof _nitroglycerin_, a slightly yellowish oil characterized by itsexplosive properties. Dynamite consists of porous earth which hasabsorbed nitroglycerin, and its strength depends on the amount present. It is used much more largely than nitroglycerin itself, since it doesnot explode so readily by concussion and hence can be transported withsafety. ~The fats. ~ These are largely mixtures of the ethereal salts knownrespectively as olein, palmitin, and stearin. These salts may beregarded as derived from oleic, palmitic, and stearic acidsrespectively, by replacing the hydrogen of the acid with the glycerinradical C_{3}H_{5}. Since this radical is trivalent and oleic, palmitic, and stearic acids contain only one replaceable hydrogen atom to themolecule, it is evident that three molecules of each acid must enterinto each molecule of the ethereal salt. The formulas for the acids andthe ethereal salts derived from each are as follows: HC_{18}H_{33}O_{2} (oleic acid) C_{8}H_{6}(C_{18}H_{33}O_{2})_{3}, (olein) HC_{16}H_{31}O_{2} (palmitic acid) C_{3}H_{5}(C_{16}H_{31}0_{2})_{3} (palmitin) HC_{18}H_{35}O_{2} (stearic acid) C_{3}H_{5}(C_{18}H_{35}O_{2})_{3} (stearin) Olein is a liquid and is the main constituent of liquid fats. Palmitinand stearin are solids. ~Butter fat and oleomargarine. ~ Butter fat consists principally of olein, palmitin, and stearin. The flavor of the fat is due to the presence of asmall amount of butyrin, which is an ethereal salt of butyric acid. Oleomargarine differs from butter mainly in the fact that a smalleramount of butyrin is present. It is made from the fats obtained fromcattle and hogs. This fat is churned up with milk, or a small amount ofbutter is added, in order to furnish sufficient butyrin to impart thebutter flavor. ~Saponification. ~ When an ethereal salt is heated with an alkali areaction expressed by the following equation takes place: C_{2}H_{5}NO_{3} + KOH = C_{2}H_{5}OH + KNO_{3}. This process is known as _saponification_, since it is the one whichtakes place in the manufacture of soaps. The ordinary soaps are made byheating fats with a solution of sodium hydroxide. The reactionsinvolved may be illustrated by the following equation representing thereaction between palmitin and sodium hydroxide: C_{3}H_{5}(C_{16}H_{31}O_{2})_{3} + 3 NaOH = 3 NaC_{16}H_{31}O_{2} + C_{3}H_{5}(OH)_{3}. In accordance with this equation the ethereal salts in the fats areconverted into glycerin and the sodium salts of the corresponding acids. The sodium salts are separated and constitute the soaps. These salts aresoluble in water. When added to water containing calcium salts theinsoluble calcium palmitate and stearate are precipitated. Magnesiumsalts act in a similar way. It is because of these facts that soap isused up by hard waters. ETHERS When ethyl alcohol is heated to 140° with sulphuric acid the reactionexpressed by the following equation takes place: 2C_{2}H_{5}OH = (C_{2}H_{5})_{2}O + H_{2}O. The resulting compound, (C_{2}H_{5})_{2}O, is ordinary ether and is themost important member of the class of compounds called _ethers_. Ordinarily ether is a light, very inflammable liquid boiling at 35°. Itis used as a solvent for organic substances and as an anęsthetic insurgical operations. KETONES The most common member of this group is acetone (C_{3}H_{6}O), acolorless liquid obtained when wood is heated in the absence of air. Itis used in the preparation of other organic compounds, especiallychloroform. ORGANIC BASES This group includes a number of compounds, all of which contain nitrogenas well as carbon. They are characterized by combining directly withacids to form salts, and in this respect they resemble ammonia. Theymay, indeed, be regarded as derived from ammonia by displacing a part orall of the hydrogen present in ammonia by hydrocarbon radicals. Amongthe simplest of these compounds may be mentioned methylamine(CH_{3}NH_{2}) and ethylamine (C_{2}H_{5}NH_{2}). These two compoundsare gases and are formed in the distillation of wood and bones. Pyridine(C_{5}H_{6}N) and quinoline (C_{9}H_{7}N) are liquids present in smallamounts in coal tar, and also in the liquid obtained by the distillationof bones. Most of the compounds now classified under the general name of_alkaloids_ (which see) also belong to this group. CARBOHYDRATES The term "carbohydrate" is applied to a class of compounds whichincludes the sugars, starch, and allied bodies These compounds containcarbon, hydrogen, and oxygen the last two elements generally beingpresent in the proportion in which they combine to form water. The mostimportant members of this class are the following: Cane sugar C_{12}H_{22}O_{11}. Milk sugar C_{12}H_{22}O_{11}. Dextrose C_{6}H_{12}O_{6}. Levulose C_{6}H_{12}O_{6}. Cellulose C_{6}H_{10}O_{5}. Starch C_{6}H_{10}0_{5}. ~Cane sugar~ (C_{12}H_{22}O_{11}). This is the well-known substancecommonly called sugar. It occurs in many plants especially in the sugarcane and sugar beet. It was formerly obtained almost entirely from thesugar cane, but at present the greatest amount of it comes from thesugar beet. The juice from the cane or beet contains the sugar insolution along with many impurities. These impurities are removed, andthe resulting solution is then evaporated until the sugar crystallizesout. The evaporation is conducted in closed vessels from which the airis partially exhausted. In this way the boiling point of the solution islowered and the charring of the sugar is prevented. It is impossible toremove all the sugar from the solution. In preparing sugar from sugarcane the liquors left after separating as much of it as possible fromthe juice of the cane constitute ordinary molasses. Maple sugar is madeby the evaporation of the sap obtained from a species of the maple tree. Its sweetness is due to the presence of cane sugar, other productspresent in the maple sap imparting the distinctive flavor. When a solution of cane sugar is heated with hydrochloric or otherdilute mineral acid, two compounds, dextrose and levulose, are formed inaccordance with the following equation: C_{12}H_{22}O_{11} + H_{2}O = C_{6}H_{12}O_{6} + C_{6}H_{12}O_{6}. This same change is brought about by the action of an enzyme present inthe yeast plant. When yeast is added to a solution of cane sugarfermentation is set up. The cane sugar, however, does not fermentdirectly: the enzyme in the yeast first transforms the sugar intodextrose and levulose, and these sugars then undergo alcoholicfermentation. When heated to 160° cane sugar melts; if the temperature is increased toabout 215°, a partial decomposition takes place and a brown substanceknown as caramel forms. This is used largely as a coloring matter. ~Milk sugar~ (C_{12}H_{22}O_{11}). This sugar is present in the milk ofall mammals. The average composition of cow's milk is as follows: Water 87. 17% Casein (nitrogenous matter) 3. 56 Butter fat 3. 64 Milk sugar 4. 88 Mineral matter 0. 75 When _rennin_, an enzyme obtained from the stomach of calves, is addedto milk, the casein separates and is used in the manufacture of cheese. The remaining liquid contains the milk sugar which separates onevaporation; it resembles cane sugar in appearance but is not so sweetor soluble. The souring of milk is due to the fact that the milk sugarpresent undergoes _lactic fermentation_ in accordance with the equation C_{12}H_{22}O_{11} + H_{2}O = 4C_{3}H_{6}O_{3}. The lactic acid formed causes the separation of the casein, thus givingthe well-known appearance of sour milk. ~Isomeric compounds. ~ It will be observed that cane sugar and milk sugarhave the same formulas. Their difference in properties is due to thedifferent arrangement of the atoms in the molecule. Such compounds aresaid to be isomeric. Dextrose and levulose are also isomeric. ~Dextrose~ (_grape sugar, glucose_) (C_{6}H_{12}O_{6}). This sugar ispresent in many fruits and is commonly called grape sugar because of itspresence in grape juice. It can be obtained by heating cane sugar withdilute acids, as explained above; also by heating starch with diluteacids, the change being as follows: C_{6}H_{10}6_{5} + H_{2}O = C_{6}H_{12}O_{6}. Pure dextrose is a white crystalline solid, readily soluble in water, and is not so sweet as cane sugar. In the presence of yeast it undergoesalcoholic fermentation. It is prepared from starch in large quantities, and being less expensive than cane sugar, is used as a substitute for itin the manufacture of jellies, jams, molasses, candy, and other sweets. The product commonly sold under the name of _glucose_ contains about 45%of dextrose. ~Levulose~ _(fruit sugar)_(C_{6}H_{12}O_{6}). This sugar is a white solidwhich occurs along with dextrose in fruits and honey. It undergoesalcoholic fermentation in the presence of yeast. ~Cellulose~ (C_{6}H_{10}O_{5}). This forms the basis of all woody fibers. Cotton and linen are nearly pure cellulose. It is insoluble in water, alcohol, and dilute acids. Sulphuric acid slowly converts it intodextrose. Nitric acid forms nitrates similar to nitroglycerin incomposition and explosive properties. These nitrates are variously knownas nitrocellulose, pyroxylin, and gun cotton. When exploded they yieldonly colorless gases; hence they are used especially in the manufactureof smokeless gunpowder. _Collodion_ is a solution of nitrocellulose in amixture of alcohol and ether. _Celluloid_ is a mixture of nitrocelluloseand camphor. _Paper_ consists mainly of cellulose, the finer gradesbeing made from linen and cotton rags, and the cheaper grades from strawand wood. ~Starch~ (C_{6}H_{10}O_{5}). This is by far the most abundant carbohydratefound in nature, being present especially in seeds and tubers. In theUnited States it is obtained chiefly from corn, nearly 80% of which isstarch. In Europe it is obtained principally from the potato. Itconsists of minute granules and is practically insoluble in cold water. These granules differ somewhat in appearance, according to the source ofthe starch, so that it is often possible to determine from what plantthe starch was obtained. When heated with water the granules burst andthe starch partially dissolves. Dilute acids, as well as certainenzymes, convert it into dextrose or similar sugars. When seedsgerminate the starch present is converted into soluble sugars, which areused as food for the growing plant. ~Chemical changes in bread making. ~ The average composition of wheat flouris as follows: Water. 13. 8% Protein (nitrogenous matter) 7. 9 Fats 1. 4 Starch 76. 4 Mineral matter 0. 5 In making bread the flour is mixed with water and yeast, and theresulting dough set aside in a warm place for a few hours. The yeastfirst converts a portion of the starch into dextrose or a similar sugar, which then undergoes alcoholic fermentation. The carbon dioxide formedescapes through the dough, making it light and porous. The yeast plantthrives best at about 30°; hence the necessity for having the dough in awarm place. If the temperature rises above 50°, the vitality of theyeast is destroyed and fermentation ceases. In baking the bread, theheat expels the alcohol and also expands the bubbles of carbon dioxidecaught in the dough, thus increasing its lightness. SOME DERIVATIVES OF BENZENE Attention has been called to the complex nature of coal tar. Among thecompounds present are the hydrocarbons, benzene, toluene, naphthalene, and anthracene. These compounds are not only useful in themselves butserve for the preparation of many other important compounds known underthe general name of coal-tar products. ~Nitrobenzene~ (_oil of myrbane_) (C_{6}H_{5}NO_{2}). When benzene istreated with nitric acid a reaction takes place which is expressed bythe following equation: C_{6}H_{6} + HNO_{3} = C_{6}H_{5}NO_{2} + H_{2}O. The product C_{6}H_{5}NO_{2} is called nitrobenzene. It is a slightlyyellowish poisonous liquid, with a characteristic odor. Its main use isin the manufacture of aniline. ~Aniline~ (C_{6}H_{5}NH_{2}). When nitrobenzene is heated with iron andhydrochloric acid the hydrogen evolved by the action of the iron uponthe acid reduces the nitrobenzene in accordance with the followingequation: C_{6}H_{5}NO_{2} + 6H = C_{6}H_{5}NH_{2} + 2H_{2}O. The resulting compound is known as aniline, a liquid boiling at 182°. When first prepared it is colorless, but darkens on standing. Largequantities of it are used in the manufacture of the _aniline or coal-tardyes_, which include many important compounds. ~Carbolic acid~ (C_{6}H_{5}OH). This compound, sometimes known as_phenol_, occurs in coal tar, and is also prepared from benzene. Itforms colorless crystals which are very soluble in water. It is stronglycorrosive and very poisonous. ~Naphthalene and anthracene. ~ These are hydrocarbons occurring along withbenzene in coal tar. They are white solids, insoluble in water. Thewell-known _moth balls_ are made of naphthalene. Large quantities ofnaphthalene are used in the preparation of _indigo_, a dye formerlyobtained from the indigo plant, but now largely prepared by laboratorymethods. Similarly anthracene is used in the preparation of the dye_alizarin_, which was formerly obtained from the madder root. THE ALKALOIDS This term is applied to a group of compounds found in many plants andtrees. They all contain nitrogen, and most of them are characterized bytheir power to combine with acids to form salts. This property isindicated by the name alkaloids, which signifies alkali-like. The saltsare soluble in water, and on this account are more largely used than thefree alkaloids, which are insoluble in water. Many of the alkaloids areused in medicine, some of the more important ones being given below. ~Quinine. ~ This alkaloid occurs along with a number of others in the barkof certain trees which grow in districts in South America and also inJava and other tropical islands. It is a white solid, and its sulphateis used in medicine in the treatment of fevers. ~Morphine. ~ When incisions are made in the unripe capsules of one of thevarieties of the poppy plant, a milky juice exudes which soon thickens. This is removed and partially dried. The resulting substance is theordinary _opium_ which contains a number of alkaloids, the principal onebeing morphine. This alkaloid is a white solid and is of great servicein medicine. Among the other alkaloids may be mentioned the following: _Nicotine_, avery poisonous liquid, the salts of which occur in the leaves of thetobacco plant; _cocaine_, a crystalline solid present in coca leaves andused in medicine as a local anęsthetic; _atropine_, a solid present inthe berry of the deadly nightshade, and used in the treatment ofdiseases of the eye; _strychnine_, a white, intensely poisonous solidpresent in the seeds of the members of the _Strychnos_ family. INDEX Acetaldehyde 405 Acetic acid 406 Acetone 411 Acetylene 203 series 399 Acids 106 binary 113 characteristics 106 definition 107 dibasic 159 familiar 106 monobasic 159 nomenclature 113 organic 405 preparation 141 strength 111 ternary 113 undissociated 107 Acker furnace, 279 Agate 260 Air 83 a mechanical mixture 89 carbon dioxide in 87 changes in composition 87 liquid 91 nitrogen in 87 oxygen in 85 poisonous effects of exhaled 88 properties 90 quantitative analysis of 85 regarded as an element 83 standard for density 229 water vapor in 87 Alabaster 308 Alchemists 9 Alchemy 9 Alcohol, common 402 denatured 404 ethyl 402 methyl 402 wood 402 Alcohols 401 Aldehydes 405 Alizarin 418 Alkali 107, 274 family 274 Alkaline-earth family 300 Alkaloids 418 Allotropic forms 22 Alloys 252 Alum 333 ammonium 334 ammonium chrome 384 ammonium iron 352 baking powders 335 potassium 333 potassium chrome 384 potassium iron 352 Aluminates 332 Aluminium 327 bronze 330, 359 chloride 333 family 327 hydroxide 332 metallurgy 328 occurrence 327 oxide 331 preparation 328 properties 329 silicates 335 uses 330 Amalgam 362 Amethyst 260, 331 Ammonia 123 composition 127 preparation 123 properties 124 uses 125 Ammonium 126 acid carbonate 295 carbonate 295 chloride 294 compounds 294 Ammonium hydrosulphide 296 hydroxide 126 molybdate 388 oxalate 407 sulphate 295 sulphide 295 sulphide, yellow 296 Analysis 40 Anhydride 135 carbonic 206 chromic 387 nitric 135 nitrous 135 phosphoric 243 sulphuric 153 Anhydrite 288 Aniline 417 Anion 106 Anode 99 Anthracene 418 Antimony 250 acids 251 alloys 253 chloride 252 metallic properties 252 occurrence 251 oxides 251 preparation 251 properties 251 sulphides 251 Apatite 175, 239, 311 Aqua ammonia 124 Aqua regia 185 Aqueous tension 25 Argon 80 Arsenic 246 acids 250 antidote 250 Marsh's test 248 occurrence 246 oxides 249 preparation 246 properties 247 sulphides 250 white 249 Arsenopyrites 246 Arsine 247 Asbestos 321, 336 Atmosphere 83 constituents 83 function of constituents 84 Atomic hypothesis 61 theory 59 and laws of matter 63 and radium 314 weights, 65 accurate determination 231 and general properties 167 and specific heats 233 calculation of 231 Dalton's method 223 direct determination 233 from molecular weights 230 relation to equivalent 224 standard for 66 steps in determining 224 Atoms 62 size 65 Atropine 419 Aurates 396 Avogadro's hypothesis 226 and chemical calculations 235 and molecular weights 227 Azote 78 Azurite 357 Babbitt metal 253 Bacteria 85 decomposition of organic matter by 122 nitrifying 85 Baking powders 285, 408 alum 335 soda 285 Barium 312 chloride 313 nitrate 313 oxides 312 sulphate 313 Barytes 312 Bases 107 characteristics 107 definition 108 familiar 107 nomenclature 113 organic 412 strength 113 undissociated 108 Basic lining process 346 Bauxite 332 Beer 404 Benzene 417 derivatives 417 series 399 Benzine 400 Bessemer process 345 Bismuth 253 basic salts 255 chloride 253 nitrate 253 occurrence 253 oxides 254 preparation 253 salts, hydrolysis of 254 subnitrate 256 uses 253 Bismuthyl chloride 256 Blast furnace 341 lamp 38 Bleaching powder 306 Bleaching by chlorine 181 by sulphurous acid 152 Boiler scale 320 Bone ash 311 Bone black 200 Borax 265 bead 266 Bornite 357 Boron 257, 264 acids 265 fluoride 264 hydride 264 occurrence 264 oxides 264 preparation 264 properties 264 Brass 323 Bread making 416 Bromides 190 Bromine 187 occurrence 187 oxygen compounds 190 preparation 187 properties 188 Bronze 359 aluminium 330, 359 Butter fat 410 Butyric acid 407 By-product 284 Cadmium 325 compounds 326 Cęsium 294 Calamine 321 Calcite 305 Calcium 301 carbide 203, 310 carbonate 305 chloride 306 fluoride 308 hydroxide 303 occurrence 301 oxide 302 phosphate 246, 311 preparation 302 sulphate 308 Calomel 363 Calorie 76 Caramel 414 Carbohydrates 413 Carbolic acid 417 Carbon 196 allotropic forms 196 amorphous 198 compounds 196 crystalline forms 197 cycle in nature 88 dioxide 204 and bases 206 and plant life 88 in air 87 occurrence 204 preparation 204 properties 204 solid 204 disulphide 160, 210 family 196 hydrogen compounds 201 monoxide 208 occurrence 196 oxides 203 properties 200 pure 198 retort 199 uses 200 Carbonates 207 acid 207 Carbonic acid 206 Carborundum 259 Carnallite 288 Casein 414 Cassiterite 370 Catalysis 153 Catalyzers 153 Cathode 99 Cation 106 Caustic potash 288 soda 278 Celestite 312 Celluloid 415 Cellulose 415 Cement 304 Ceramic industries 336 Cerium 377 Chalcedony 260 Chalcocite 357 Chalcopyrite 357 Chalk 305 Chamber acid 157 Changes, physical and chemical 2 Charcoal 199 Chemical affinity 12 changes 2 compounds 7 equilibrium 128 properties 3 Chemistry, definition 4 Chili saltpeter 191, 285 Chinaware 336 Chloric acid 187 Chlorides 186 Chlorine 177 bleaching action 181 chemical properties 180 family 174 historical 177 occurrence 178 oxides 187 oxygen acids 187 preparation 178 properties 179 Chloroform 401 Chloroplatinic acid 393 Chlorous acid 187 Chromates 385 Chrome alum 384 Chromic acid 388 anhydride 387 chloride 383 hydroxide 383 sulphate 384 sulphide 384 Chromite 383 Chromium 383 a base-forming element 383 an acid-forming element 385 occurrence 383 Cinnabar 363 Citric acid 408 Clay 336 Coal 199 gas 217 products 400 tar 218 Cobalt 354 compounds 354 Cocaine 419 Coke 199 Collodion 415 Colemanite 265 Combining weights 225 Combustion 17 broad sense 20 in air 19 phlogiston theory 19 products 18 spontaneous 20 supporters 213 Compounds, chemical 7 isomeric 414 of metals, preparation 265 structure of 118 Conservation of energy 4 of matter 5 Contact process 154 Converter, Bessemer 345 Copper 357 acetate 407 alloys of 359 family 356 hydroxide 360 metallurgy 357 occurrence 357 ores 357 oxide 360 properties 358 refining 358 sulphate 361 sulphide 361 uses 359 Copperas 350 Coral 305 Corrosive sublimate 363 Corundum 331 Cream of tartar 408 Crocoisite 383 Cryolite 175, 328 Crystallization 98 water of 54, 75 Crystallography 161 Crystals 161 axes of 161 systems 162 Cupric compounds 360 Cuprite 360 Cuprous compounds 360 chloride 360 oxide 360 Cyanides 210 solutions are alkaline 210 Dalton's atomic hypothesis 61 Decay 21 Decomposition of organic matter 122 Decrepitation 55 Deliquescence 55 Density of gases 230 Desiccating agents 55 Developers 367 Dewar bulb 91 Dextrose 414 Diamond 197 Dichromates 385 Dichromic acid 385 Dimorphous substances 163 Dissociation 99 and boiling point 101 and freezing point 101 equations of 112 extent of 113 Distillation 50 Dogtooth spar 306 Dolomite 319 Double decomposition 71 Drummond light 38 Dyeing 333 Dynamite 409 Earth metals 327 Efflorescence 54 Electric furnace 221 Electro-chemical industries 269 Electrode 99 Electrolysis 99 of sodium chloride 102 of sodium sulphate 103 of water 41, 102 Electrolytes 99 Electrolytic dissociation 99 Electroplating 366 Electrotyping 359 Elements, definition 8 atomic weights 232 earlier classification 165 names 11 natural groups 165 number of 9 occurrence 10 periodic division 166 physical state 10 symbols of 11 Emery 331 Energy 4 and plant life 89 chemical 5 conservation of 4 transformation of 5 Enzyme 405 Epsom salts 320 Equations 68 are quantitative 72 knowledge requisite for 69 not algebraic 74 reading of 69 Equilibrium 138 chemical 138 in solution 139 point of 138 Equivalent 224 determination of 224 elements with more than one 225 relation to atomic weight 224 Etching 177 Ether 411 Ethereal salts 409 Ethers 411 Ethylamine 412 Ethylene series 399 Eudiometer 43 Evaporation 11 Families in periodic groups 170 triads 165 Family resemblances 170 Fats 409 Fatty acid series 405 Feldspar 261, 335 Fermentation 404 acetic 406 alcoholic 404, 405 lactic 414 Ferric chloride 352 hydroxide 352 salts 351 reduction 353 sulphate 352 Ferrochromium, 383 Ferromanganese 343 Ferrosilicon 259 Ferrous carbonate 351 salts 350 oxidation of 353 sulphate 350 sulphide 350 Fertilizers 245 Filtration 6, 51 beds 52 Fire damp 202 Flames 213 appearance 214 blowpipe 216 Bunsen 214 conditions for 213 hydrogen 34 luminosity 216 oxidizing 214 oxyhydrogen 37 reactions 296 reducing 214 structure 214 Flash lights 317 Flint 260 Fluorides 177 Fluorine 175 Fluorspar 175, 308 Fluosilicic acid 259 Flux 340 Fool's gold 351 Formaldehyde 405 Formalin 405 Formic acid 406 Formulas 68 how determined 234 structural 119 Fractional distillation 51 Franklinite 321 Fuels 220 Furnace, arc 221 electric 221 resistance 221 Fusion methods 271 Galena 373 Gallium 327 Galvanized iron 323 Gas, collection of 15 coal 217 fuel 217 illuminating 217 measurement of 23 natural 219 purification of 218 water 219 Gases, table 220 Gasoline 400 German silver 323, 359 Germanium 370 Germs, effect of cold on 53 in air 84 in water 52 Glass 262 coloring of 263 etching of 177 molding of 263 nature of 263 varieties 263 Glauber's salt 281 Glazing 336 Glucose 414 Glycerin 405 nitrates of 409 Gold 393 alloys 396 chloride 396 coin 359 extraction of 394 in copper 358 mining 394 occurrence 393 properties 396 refining of 395 telluride 394 Goldschmidt method 269, 330 Gram-molecular weight 236 Granite 336 Graphite 198 Gun cotton 415 metal 359 powder 292 Gypsite 308 Gypsum 308 Halogens 174 Hard water 309 Heat of reaction 75 Helium 80, 314 Hematite 339, 349 Homologous series 398 Hydriodic acid 193 Hydrobromic acid 189 Hydrocarbons 201, 398 properties 400 series 398 substitution products 401 Hydrochloric acid 182 composition 183 oxidation of 185 preparation 182 properties 184 salts 186 Hydrocyanic acid 210 Hydrofluoric acid 176 etching by 177 salts of 177 Hydrogen 28 dioxide 56 explosive with oxygen 35 occurrence 28 preparation from acids 30 preparation from water 28 properties 32 standard for atomic weights 66 standard for molecular weights 227 sulphide 146 uses 38 Hydrolysis 254 conditions affecting 255 partial 255 Hydrosulphuric acid 146 Hydroxyl radical 112 Hypochlorous acid 187 Hypothesis 61 Avogadro's 226 Dalton's 61 Ice manufacture 125 Iceland spar 305 Indigo 418 Indium 327 Insoluble compounds 272 Iodic acid 194 Iodides 193 Iodine 190 oxygen compounds 193 preparation 191 properties 192 tincture 192 Iodoform 192, 401 Ions 100 and electrolytes 104 Iridium 393 Iron 339 alum 352 cast 343 compounds 349 cyanides 352 disulphide 351 family 338 metallurgy 339 occurrence 339 ores 339 oxides 349 pure 348 varieties 342, 347 wrought 343 Jasper 260 Kainite 288 Kaolin 261, 335 Kerosene 400 Ketones 411 Kieserite 288 Kindling temperature 17 Krypton 80 Lactic acid 408 Lampblack 200 Laughing gas 132 Law, definition 61 of Boyle 24 of Charles 23 of combining volumes 194 of conservation of energy 4 of conservation of matter 5, 59 of definite composition 59 of Dulong and Petit 233 of Gay-Lussac 194 of multiple proportion 60 of Raoult 233 periodic 169 Lead 373 acetate 375, 407 alloys 375 basic carbonate 376 carbonate 376 chloride 377 chromate 377 insoluble compounds 376 metallurgy 373 nitrate 375 occurrence 373 oxides 375 peroxide 375 properties 374 red 375 soluble salts 375 sugar of 375 sulphate 377 sulphide 377 white 376 Le Blanc soda process 282 Levulose 415 Lime 302 air-slaked 303 hypochlorite 307 kilns 303 slaked 303 Lime light 38 Limestone 305 Limewater 303 Limonite 339 Litharge 375 Lithium 294 Luminosity of flames 216 Lunar caustic 366 Magnesia 318 alba 319 usta 318 Magnesite 318 Magnesium 317 basic carbonate 319 carbonate 318 cement 318 chloride 319 family 316 hydroxide 318 oxide 318 silicates 321 sulphate 320 Magnetite 339, 349 Malachite 357 Manganates 381 Manganese 379 a base-forming element 380 an acid-forming element 381 in glass 263 occurrence 379 oxides 380 Manganic acid 381 Manganous salts 380 Marble 305 Marl 305 Marsh gas 202 Matches 242 Matte 358 Matter, classification 6 conservation 5 definition 5 kinds 9 Measurement of gases 23 Mechanical mixtures 6 Meerschaum 321, 336 Mercuric chloride 363 iodide 364 oxide 14, 362 sulphide 363 Mercurous chloride 363 Mercury 361 iodides 364 metallurgy 361 occurrence 361 oxides 362 uses 362 Metaboric acid 265 Metallurgy 268 Metals 165, 267 action on salts 271 definition 267 extraction 268 occurrence 267 preparation of compounds 269 reduction from ores 268 Metaphosphoric acid 245 Metarsenic acid 250 Metasilicic acid 261 Metastannic acid 371 Methane 202, 399 Methylamine 412 Mexican onyx 305 Mica 261, 336 Microcosmic salt 244 Milk 414 Minerals 267 Minium 375 Mixed salts 244 Molasses 413 Molecular weights 226 boiling-point method 233 compared with oxygen 228 determination 226 freezing-point method 233 oxygen standard 227 of elements 232 vapor-density method 229 Molecule 62 Molybdenum 388 Molybdic acid 388 Monazite sand 377 Mordants 333 Morphine 418 Mortar 304 Moth balls 418 Muriatic acid 182 Naphthalene 418 Naphthas 400 Nascent state 182 Natural gas 219 sciences 1 Neon 80 Neutralization 108 a definite act 109 definition 109 heat of 109 partial 111 Niagara Falls 269, 329 Nickel 354 coin 359 compounds 354 plating 354 Nicotine 419 Nitrates 131 Nitric acid, 128 action on metals 130 decomposition 129 oxidizing action 130 preparation 128, 140 properties 129 salts 131 Nitric oxide 133 Nitrites 132 Nitrobenzene 417 Nitrocellulose 415 Nitrogen 78 compounds 122 in air 87 occurrence 78, 122 oxides 132 preparation 78 properties 80 Nitroglycerin 409 Nitrosulphuric acid 155 Nitrous acid 132 oxide 132 Non-metals 165 Oil of myrbane 417 of vitriol 154 Oleic acid 408 Olein 409 Oleomargarine 410 Onyx 260 Opal 260 Open-hearth process 346 Opium 418 Ores 267 Organic bases 412 chemistry 201, 397 matter, decomposition 122 Orpiment 246 Orthoarsenic acid 250 Orthophosphates 244 Orthophosphoric acid 244 Orthosilicic acid 261 Osmic acid 393 Osmium 393 tetroxide 393 Oxalic acid 407 Oxidation 17, 353 definition 18 Oxidizing agent 37 Oxygen 13 and ozone 22 commercial preparation 16 history 13 importance 21 in air estimation, 85 in air function, 84 occurrence 13 preparation 13 properties 16 standard for atomic weights 66 two atoms in molecule 227 Oxyhydrogen blowpipe 37 Ozone 21, 137 Palladium 390 Palmitic acid 407 Palmitin 409 Paraffin 400 Paris green 250 Parkes's method for silver 364 Pearls 305 Perchloric acid 187 Periodic acid 194 Periodic division 166 groups 167 law 169 law, imperfections 172 law, value 171 table 168 table, arrangement 166 Permanent hardness 310 Permanganates 381 Permanganic acid 381 Peroxides 278 Petroleum 399 Pewter 372 Phenol 417 Philosopher's stone 9 Phlogiston 19 Phosphates 245 Phosphine 242 Phosphonium compounds 243 Phosphoric acid 244 Phosphorite 239 Phosphorous acid 244 Phosphorus 239 acids 243 family 238 hydrogen compounds 242 occurrence 239 oxides 243 preparation 239 properties 240 red 241 yellow 240 Photography 367 Physical changes 2 properties 3 properties and periodic groups 171 state 3 Physics 1, 4 Pitchblende 314 Plaster of Paris 308 Platinic chloride 393 Platinized asbestos 391 Platinous chloride 393 Platinum 391 a catalytic agent 152, 392 Pneumatic trough 16 Polyboric acid 265 Polyhalite 288 Polysilicic acids 261 Porcelain 336 Portland cement 304 Potash 293 Potassium 287 acid carbonate 294 acid sulphate 294 acid sulphite 294 alum, aluminium 334 alum, chrome 384 alum, iron 352 and plant life 287 aurate 396 bromide 290 carbonate 293 chlorate 291 chloride 290 chromate 385 cyanide 293 dichromate 386 ferricyanide 352 ferrocyanide 352 hydroxide 288 hydroxide, action of halogens 289 hypochlorite 289 iodide 290 manganate 381 nitrate 291 occurrence 287 permanganate 381 preparation 288 sulphate 294 Precipitated chalk 306 Precipitation 140 Properties, chemical 3 physical 3 Prussic acid 210 Puddling 343 furnace 344 Pyridine 412 Pyrites 351 Pyrolusite 380 Pyrophosphoric acid 245 Quantitative equations 72 Quartz 260 Quicklime 302 Quinine 418 Quinoline 412 Radical 112 Radium 313 Reaction, classes 70 addition 70 completed 139 heat of 75 of decomposition 70 of double decomposition 71 of substitution 70 reversible 137 steps in 131 Realgar 246 Red lead 375 phosphorus 241 Reducing agent 37 Reduction 36, 354 Rennin 414 Resemblances, family 170 Respiration 87 Rhodium 390 Rochelle salts 408 Rouge 349 Rubidium 294 Ruby 331 Ruthenium 390 Rutile 264 Safety lamp 202 Sal ammoniac 294 soda 282 Salt 280 Saltpeter 291 Chili 285 Salts, 109 acid, 112 Salts basic 111 binary 114 characteristics 109 definition 109 insoluble 272 mixed 244 nomenclature 113 normal 112 preparation by precipitation 270 Sand 260 Sandstone 260 Saponification 410 Sapphire 331 Satinspar 308 Scale 320 Schönite 288 Selenite 308 Selenium 161 Serpentine 320, 336 Shot 247, 375 Siderite 339 Silica 260 Silicates 261 Silicic acids 261 Silicides 259 Silicon 258 acids 261 dioxide 260 fluoride 258 hydride 258 Silver 364 amalgamation process 364 bromide 367 chloride 367 coin 359 German 359 in copper ores 358 iodide 367 metallurgy 364 nitrate 366 oxide 366 parting of 365 refining 365 sulphide 366 Slag 340 Smalt 355 Smithsonite 321 Smokeless powder 293 Soaps 410 Soda ash 284 Soda lime 202 Sodium 276 acetate 407 bicarbonate 285 carbonate 282 carbonate, historical 284 chloride 280 chromates 386 hydrogen carbonate 285 hydroxide 278 hyposulphite 282 iodate 191 nitrate 285 occurrence 276 peroxide 277 phosphates 286 preparation 276 properties 277 sulphate 281 sulphite 281 tetraborate 287 thiosulphate 282 Solder 372, 375 Solubility of gases 95 of solids 96 Solution 94 and chemical action 53 boiling point 98 classes 94 distribution of solids in 98 electrolysis of 99 freezing point 99 of gases in liquids 94 of solids in liquids 96 properties 98 saturated 97 supersaturated 98 Solvay soda process 283 Sombrerite 239 Spectroscope 296 Sphalerite 325 Spiegel iron 343 Spinel 332 Spontaneous combustion 20 Stalactites 305 Stalagmites 305 Standard conditions 23 Stannates 372 Stannic acid 372 chloride 372 oxide 372 Stannous chloride 372 Starch 415 Stassfurt salts 287 Stearic acid 407 Stearin 409 Steel 345 alloys 348 properties 347 tempering of 348 tool 347 Stibine 251 Stibnite 250 Stoneware 336 Strontianite 312 Strontium 312 hydroxide 312 nitrate 312 Structural formulas 119 Structure of compounds 119 Strychnine 419 Substitution 70 Sugars 412 cane 412 fruit 415 grape 414 milk 414 Sulphates 159 Sulphides 148 Sulphites 152 action of acids on 150 Sulphur 143 allotropic forms 144 chemical properties 145 comparison with oxygen 161 dioxide 149 preparation 149 properties 150 extraction 143 flowers of 143 occurrence 143 oxides 149 physical properties 144 trioxide 152 uses 146 varieties 144 Sulphuric acid 154 action as an acid 157 action on metals 157 action on organic matter 158 action on salts 158 action on water 158 fuming 155 manufacture 154 oxidizing action 157 plant 156 properties 157 salts 159 Sulphuric anhydride 153 Sulphurous acid 151 Superphosphate of lime 246 Sylvine 288 Symbols 11 Synthesis 40 Table, alkali metals 274 alkaline-earth metals 300 alloys of copper 359 aqueous tension Appendix B atomic weights Appendix A chlorine family 174 composition of earth's crust 10 composition of fuel gases 220 constants of elements Appendix B copper family 356 elements Appendix A gold and platinum metals 390 hydrocarbons 399 magnesium family 316 manganese and chromium 379 periodic arrangement 168 phosphorus family 238 silicon family 257 solubility of gases in water 95 solubility of salts 96 solubility of salts at different temperatures 97 tin and lead 370 weights of gases Appendix B Talc 321, 336 Tartar emetic 408 Tartaric acid 408 Tellurium 161 Temporary hardness 309 Ternary acids 113 salts 114 Tetraboric acid 265 Thallium 327 Theory, atomic 61 definition 64 value of 64 Thermite 331 Thio compounds 282 Thiosulphates 159 Thiosulphuric acid 159 Thorium 377 Tin 370 block 371 compounds 372 crystals 372 family 370 foil 371 metallurgy 370 plate 371 properties 371 uses 371 Titanium 257, 264 Topaz 331 Triad families 166 Tungsten 388 Type metal 253, 375 Uranium 388 Valence 116 a numerical property 116 and combining ratios 118 and equations 120 and formulas 120 and periodic groups 162 and structure 118 definition 116 indirectly determined 117 measure of 117 variable 117 Vaseline 400 Venetian red 349 Verdigris 407 Vermilion 363 Vinegar 406 Vitriol, blue 361 green 350 oil of 154 white 324 Volume and aqueous tension 25 and pressure 24 and temperature 23 of combining gases 194 Water 40 a compound 40 and disease 49 catalytic action of 154 chalybeate 351 chemical properties 53 composition 47 composition by volume 44 composition by weight 47 dissociation of 210 distillation of 50 electrolysis of 41, 103 filtration of 51 gas 219 hard 309 historical 40 impurities in 48 in air 87 mineral 49 occurrence 48 of crystallization 54, 75 physical properties 53 purification of 50 qualitative analysis 41 quantitative analysis 42 river 49 sanitary analysis 50 self-purification 53 softening of 310 standard substance 55 synthesis 43 uses of 55 Weights, atomic 65 Welsbach mantles 219, 377 Whisky 404 Wine 404 Witherite 312 Wood alcohol 402 distillation 402 Wood's metal 254 Xenon 80 Yeast 403 Zinc 321 alloys of 323 blende 321 chloride 325 flowers of 322 metallurgy 321 occurrence 321 oxide 324 sulphate 324 sulphide 325 white 324 Zymase, 403 ANNOUNCEMENTS AN ELEMENTARY STUDY OF CHEMISTRY By WILLIAM McPHERSON, Professor of Chemistry in Ohio State University, and WILLIAM E. HENDERSON, Associate Professor of Chemistry in Ohio StateUniversity. 12mo. Cloth. 434 pages. Illustrated. List price, $1. 25; mailing price, $1. 40 This book is the outgrowth of many years of experience in the teachingof elementary chemistry. In its preparation the authors have steadfastlykept in mind the limitations of the student to whom chemistry is a newscience. They have endeavored to present the subject in a clear, well-graded way, passing in a natural and logical manner from principleswhich are readily understood to those which are more difficult to grasp. The language is simple and as free as possible from unusual andtechnical phrases. Those which are unavoidable are carefully defined. The outline is made very plain, and the paragraphing is designed to beof real assistance to the student in his reading. The book is in no way radical, either in the subject-matter selected orin the method of treatment. At the same time it is in thorough harmonywith the most recent developments in chemistry, both in respect totheory and discovery. Great care has been taken in the theoreticalportions to make the treatment simple and well within the reach of theability of an elementary student. The most recent discoveries have beentouched upon where they come within the scope of an elementary text. Especial attention has been given to the practical applications ofchemistry, and to the description of the manufacturing processes in useat the present time. EXERCISES IN CHEMISTRY. By WILLIAM McPHERSON and WILLIAM E. HENDERSON. (_In press. _) GINN & COMPANY PUBLISHERS A FIRST COURSE IN PHYSICS By ROBERT A. MILLIKAN, Associate Professor of Physics, and HENRY G. GALE, Assistant Professor of Physics in The University of Chicago 12mo, cloth, 488 pages, illustrated, $1. 25 A LABORATORY COURSE IN PHYSICS _FOR SECONDARY SCHOOLS_ By ROBERT A. MILLIKAN and HENRY G. GALE 12mo, flexible cloth, 134 pages, illustrated, 40 cents This one-year course in physics has grown out of the experience of theauthors in developing the work in physics at the School of Education ofThe University of Chicago, and in dealing with the physics instructionin affiliated high schools and academies. The book is a simple, objective presentation of the subject as opposedto a formal and mathematical one. It is intended for the third-yearhigh-school pupils and is therefore adapted in style and method oftreatment to the needs of students between the ages of fifteen andeighteen. It especially emphasizes the historical and practical aspectsof the subject and connects the study very intimately with facts ofdaily observation and experience. The authors have made a careful distinction between the class ofexperiments which are essentially laboratory problems and those whichbelong more properly to the classroom and the lecture table. The formerare grouped into a Laboratory Manual which is designed for use inconnection with the text. The two books are not, however, organicallyconnected, each being complete in itself. All the experiments included in the work have been carefully chosen withreference to their usefulness as effective classroom demonstrations. GINN AND COMPANY PUBLISHERS APPENDIX A LIST OF THE ELEMENTS, THEIR SYMBOLS, AND ATOMIC WEIGHTS The more important elements are marked with an asterisk O = 16 *Antimony Sb 120. 2*Argon A 39. 9*Arsenic As 75. 0*Barium Ba 137. 4Beryllium Be 9. 1*Bismuth Bi 208. 5*Boron B 11. 0*Bromine Br 79. 96*Cadmium Cd 112. 4Cęsium Cs 132. 9*Calcium Ca 40. 1*Carbon C 12. 00Cerium Ce 140. 25*Chlorine Cl 35. 45*Chromium Cr 52. 1*Cobalt Co 59. 0Columbium Cb 94. 0*Copper Cu 63. 6Erbium Er 166. 0*Fluorine F 19. 0Gadolinium Gd 156. 0Gallium Ga 70. 0Germanium Ge 72. 5*Gold Au 197. 2Helium He 4. 0*Hydrogen H 1. 008Indium In 115. 0*Iodine I 126. 97Iridium Ir 193. 0*Iron Fe 55. 9Krypton Kr 81. 8Lanthanum La 138. 9*Lead Pb 206. 9Lithium Li 7. 03*Magnesium Mg 24. 36*Manganese Mn 55. 0*Mercury Hg 200. 0Molybdenum Mo 96. 0Neodymium Nd 143. 6Neon Ne 20. 0*Nickel Ni 58. 7*Nitrogen N 14. 04Osmium Os 191. 0*Oxygen O 16. 00Palladium Pd 106. 5*Phosphorus P 31. 0*Platinum Pt 194. 8*Potassium K 39. 15Praseodymium Pr 140. 5Radium Ra 225. 0Rhodium Rh 103. 0Rubidium Rb 85. 5Ruthenium Ru 101. 7Samarium Sm 150. 3Scandium Sc 44. 1Selenium Se 79. 2*Silicon Si 28. 4*Silver Ag 107. 93*Sodium Na 23. 05*Strontium Sr 87. 6*Sulphur S 32. 06Tantalum Ta 183. 0Tellurium Te 127. 6Terbium Tb 160. 0Thallium Tl 204. 1Thorium Th 232. 5Thulium Tm 171. 0*Tin Sn 119. 0Titanium Ti 48. 1Tungsten W 184. 0Uranium U 238. 5Vanadium V 51. 2Xenon Xe 128. 0Ytterbium Yb 173. 0Yttrium Yt 89. 0*Zinc Zn 65. 4Zirconium Zr 90. 6 APPENDIX B Tension of Aqueous Vapor expressed in Millimeters of Mercury TEMPERATURE PRESSURE16 13. 517 14. 418 15. 319 16. 320 17. 421 18. 522 19. 623 20. 924 22. 225 23. 5 Weight of 1 Liter of Various Gases measured under Standard Conditions Acetylene 1. 1614Air 1. 2923Ammonia 0. 7617Carbon dioxide 1. 9641Carbon monoxide 1. 2499Chlorine 3. 1650Hydrocyanic acid 1. 2036Hydrochloric acid 1. 6275Hydrogen 0. 08984Hydrosulphuric acid 1. 5211Methane 0. 7157Nitric oxide 1. 3410Nitrogen 1. 2501Nitrous oxide 1. 9677Oxygen 1. 4285Sulphur dioxide 2. 8596 Densities and Melting Points of Some Common Elements DENSITY MELTING POINTAluminium 2. 68 640Antimony 6. 70 432Arsenic 5. 73 --Barium 3. 75 --Bismuth 9. 80 270Boron 2. 45 --Cadmium 8. 67 320Cęsium 1. 88 26. 5Calcium 1. 54 --Carbon, Diamond 3. 50 -- " Graphite 2. 15 -- " Charcoal 1. 80 --Chromium 7. 30 3000Cobalt 8. 55 1800Copper 8. 89 1084Gold 19. 30 1064Iridium 22. 42 1950Iron 7. 93 1800Lead 11. 38 327Lithium 0. 59 186Magnesium 1. 75 750Manganese 8. 01 1900Mercury 13. 596 -39. 5Nickel 8. 9 1600Osmium 22. 47 --Palladium 11. 80 1500Phosphorus 1. 80 45Platinum 21. 50 1779Potassium 0. 87 62. 5Rhodium 12. 10 --Rubidium 1. 52 38. 5Ruthenium 12. 26 --Silicon 2. 35 --Silver 10. 5 960Sodium 0. 97 97. 6Strontium 2. 50 --Sulphur 2. 00 114. 8Tin 7. 35 235Titanium 3. 50 --Zinc 7. 00 420