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Title: An Elementary Study of Chemistry

Author: William McPherson
        William Edwards Henderson

Release Date: March 18, 2007 [EBook #20848]

Language: English

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ANTOINE LAURENT LAVOISIER

Famous for his care in quantitative experiments, for demonstrating the true nature of combustion, for introducing system into the naming and grouping of chemical substances. Executed (1794) during the French Revolution because of his connection with the government

This picture is taken from a French engraving of 1799. The panel represents Lavoisier as he is being arrested in his laboratory by the Revolutionary Committee

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 & COMPANY
BOSTON * NEW YORK * CHICAGO * LONDON

COPYRIGHT, 1905, 1906, BY
WILLIAM MCPHERSON AND WILLIAM E. HENDERSON

ALL RIGHTS RESERVED


The Athenæum Press
GINN & COMPANY * PROPRIETORS * BOSTON * U.S.A.

Transcriber's note: Minor typos have been corrected.

[Pg iii]

PREFACE

In offering this book to teachers of elementary chemistry the authors lay no claim to any great originality. It has been their aim to prepare a text-book constructed along lines which have become recognized as best suited to an elementary treatment of the subject. At the same time they have made a consistent effort to make the text clear in outline, simple in style and language, conservatively modern in point of view, and thoroughly teachable.

The question as to what shall be included in an elementary text on chemistry is perhaps the most perplexing one which an author must answer. While an enthusiastic chemist with a broad understanding of the science is very apt to go beyond the capacity of the elementary student, the authors of this text, after an experience of many years, cannot help believing that the tendency has been rather in the other direction. In many texts no mention at all is made of fundamental laws of chemical action because their complete presentation is quite beyond the comprehension of the student, whereas in many cases it is possible to present the essential features of these laws in a way that will be of real assistance in the understanding of the science. For example, it is a difficult matter to deduce the law of mass action in any very simple way; yet the elementary student can readily comprehend that reactions are reversible, and that the point of equilibrium depends upon, rather simple conditions. The authors believe that it is worth while to[Pg iv] present such principles in even an elementary and partial manner because they are of great assistance to the general student, and because they make a foundation upon which the student who continues his studies to more advanced courses can securely build.

The authors have no apologies to make for the extent to which they have made use of the theory of electrolytic dissociation. It is inevitable that in any rapidly developing science there will be differences of opinion in regard to the value of certain theories. There can be no question, however, that the outline of the theory of dissociation here presented is in accord with the views of the very great majority of the chemists of the present time. Moreover, its introduction to the extent to which the authors have presented it simplifies rather than increases the difficulties with which the development of the principles of the science is attended.

The oxygen standard for atomic weights has been adopted throughout the text. The International Committee, to which is assigned the duty of yearly reporting a revised list of the atomic weights of the elements, has adopted this standard for their report, and there is no longer any authority for the older hydrogen standard. The authors do not believe that the adoption of the oxygen standard introduces any real difficulties in making perfectly clear the methods by which atomic weights are calculated.

The problems appended to the various chapters have been chosen with a view not only of fixing the principles developed in the text in the mind of the student, but also of enabling him to answer such questions as arise in his laboratory work. They are, therefore, more or less practical in character. It is not necessary that all of them should[Pg v] be solved, though with few exceptions the lists are not long. The answers to the questions are not directly given in the text as a rule, but can be inferred from the statements made. They therefore require independent thought on the part of the student.

With very few exceptions only such experiments are included in the text as cannot be easily carried out by the student. It is expected that these will be performed by the teacher at the lecture table. Directions for laboratory work by the student are published in a separate volume.

While the authors believe that the most important function of the elementary text is to develop the principles of the science, they recognize the importance of some discussion of the practical application of these principles to our everyday life. Considerable space is therefore devoted to this phase of chemistry. The teacher should supplement this discussion whenever possible by having the class visit different factories where chemical processes are employed.

Although this text is now for the first time offered to teachers of elementary chemistry, it has nevertheless been used by a number of teachers during the past three years. The present edition has been largely rewritten in the light of the criticisms offered, and we desire to express our thanks to the many teachers who have helped us in this respect, especially to Dr. William Lloyd Evans of this laboratory, a teacher of wide experience, for his continued interest and helpfulness. We also very cordially solicit correspondence with teachers who may find difficulties or inaccuracies in the text.

The authors wish to make acknowledgments for the photographs and engravings of eminent chemists from which[Pg vi] the cuts included in the text were taken; to Messrs. Elliott and Fry, London, England, for that of Ramsay; to The Macmillan Company for those of Davy and Dalton, taken from the Century Science Series; to the L. E. Knott Apparatus Company, Boston, for that of Bunsen.

THE AUTHORS

OHIO STATE UNIVERSITY

COLUMBUS, OHIO

[Pg vii]

CONTENTS

[Pg ix]

LIST OF FULL-PAGE ILLUSTRATIONS

[Pg 1]

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 a number of more limited ones for the convenience of the investigator as well 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 happens that the different sciences are very intimately related, and a thorough knowledge of any one of them involves a considerable acquaintance with several others. Thus the botanist must know something about animals as well as about plants; the student of human physiology must know something about physics as well as about the parts of the body.

Intimate relation of chemistry and physics. Physics and chemistry are two sciences related in this close way, and it is not easy to make a precise distinction between them. In a general way it may be said that they are both concerned with inanimate matter rather than with living, and more particularly with the changes which such matter[Pg 2] may be made to undergo. These changes must be considered more closely before a definition of the two sciences can be given.

Physical changes. One class of changes is not accompanied by an alteration in the composition of matter. When a lump of coal is broken the pieces do not differ from the original lump save in size. A rod of iron may be broken into pieces; it may be magnetized; it may be heated until it glows; it may be melted. In none of these changes has the composition of the iron been affected. The pieces of iron, the magnetized iron, the glowing iron, the melted iron, are just as truly iron as was the original rod. Sugar may be dissolved in water, but neither the sugar nor the water is changed in composition. The resulting liquid has the sweet taste of sugar; moreover the water may be evaporated by heating and the sugar recovered unchanged. Such changes are called physical changes.

DEFINITION: Physical changes are those which do not involve a change in the composition of the matter.

Chemical changes. Matter may undergo other changes in which its composition is altered. When a lump of coal is burned ashes and invisible gases are formed which are entirely different in composition and properties from the original coal. A rod of iron when exposed to moist air is gradually changed into rust, which is entirely different from the original iron. When sugar is heated a black substance is formed which is neither sweet nor soluble in water. Such changes are evidently quite different from the physical changes just described, for in them new substances are formed in place of the ones undergoing change. Changes of this kind are called chemical changes.[Pg 3]

DEFINITION: Chemical changes are those which involve a change in the composition of the matter.

How to distinguish between physical and chemical changes. It is not always easy to tell to which class a given change belongs, and many cases will require careful thought on the part of the student. The test question in all cases is, Has the composition of the substance been changed? Usually this can be answered by a study of the properties of the substance before and after the change, since a change in composition is attended by a change in properties. In some cases, however, only a trained observer can decide the question.

Changes in physical state. One class of physical changes should be noted with especial care, since it is likely to prove misleading. It is a familiar fact that ice is changed into water, and water into steam, by heating. Here we have three different substances,—the solid ice, the liquid water, and the gaseous steam,—the properties of which differ widely. The chemist can readily show, however, that these three bodies have exactly the same composition, being composed of the same substances in the same proportion. Hence the change from one of these substances into another is a physical change. Many other substances may, under suitable conditions, be changed from solids into liquids, or from liquids into gases, without change in composition. Thus butter and wax will melt when heated; alcohol and gasoline will evaporate when exposed to the air. The three states—solid, liquid, and gas—are called the three physical states of matter.

Physical and chemical properties. Many properties of a substance can be noted without causing the substance to undergo chemical change, and are therefore called its physical properties. Among these are its physical state, color, odor, taste, size, shape, weight. Other properties are only[Pg 4] discovered when the substance undergoes chemical change. These are called 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 a general distinction between physics and chemistry.

DEFINITION: Physics is the science which deals with those changes in matter which do not involve a change in composition.

DEFINITION: Chemistry is the science which deals with those changes in matter 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 do work. Thus water falling from a height upon a water wheel turns the wheel and in this way does the work of the mills. Magnetized iron attracts iron to itself and the motion of the iron as it moves towards the magnet can be made to do work. When coal is burned it causes the engine to move and transports the loaded cars from place to place. When a body has this power to do work it is said to possess energy.

Law of conservation of energy. Careful experiments have shown that when one body parts with its energy the energy is not destroyed but is transferred to another body or system of bodies. Just as energy cannot be destroyed, neither can it be created. If one body gains a certain amount of energy, some other body has lost an equivalent amount.[Pg 5] These facts are summed up in the law of conservation of energy which may be stated thus: While energy can be changed from one form into another, it cannot be created or destroyed.

Transformations of energy. Although energy can neither be created nor destroyed, it is evident that it may assume many different forms. Thus the falling water may turn the electric generator and produce a current of electricity. The energy lost by the falling water is thus transformed into the energy of the electric current. This in turn may be changed into the energy of motion, as when the current is used for propelling the cars, or into the energy of heat and light, as when it is used for heating and lighting the cars. Again, the energy of coal may be converted into energy of heat and subsequently of motion, as when it is used as a fuel in steam engines.

Since the energy possessed by coal only becomes available when the coal is made to undergo a chemical change, it is sometimes called chemical energy. It is this form of energy in which we are especially interested in the study of chemistry.

Matter. Matter may be defined as that which occupies space and possesses weight. Like energy, matter may be changed oftentimes from one form into another; and since in these transformations all the other physical properties of a substance save weight are likely to change, the inquiry arises, Does the weight also change? Much careful experimenting has shown that it does not. The weight of the products formed in any change in matter always equals the weight of the substances undergoing change.

Law of conservation of matter. The important truth just stated is frequently referred to as the law of conservation[Pg 6] of matter, and this law may be briefly stated thus: Matter can neither be created nor destroyed, though it can be changed from one form into another.

Classification of matter. At first sight there appears to be no limit to the varieties of matter of which the world is made. For convenience in study we may classify all these varieties under three heads, namely, mechanical mixtures, chemical compounds, and elements.

Fig. 1

Mechanical mixtures. If equal bulks of common salt and iron filings are thoroughly mixed together, a product is obtained which, judging by its appearance, 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 of which substances retains its own peculiar properties. The mixture tastes just like salt; the iron particles can be seen and their gritty character detected. A magnet rubbed in the mixture draws out the iron just as if the salt were not there. On the other hand, the salt can be separated from the iron quite easily. Thus, if several grams of the mixture are placed in a test tube, and the tube half filled with water and thoroughly shaken, the salt dissolves in the water. The iron particles can then be filtered from the liquid by pouring the entire mixture upon a piece of filter paper folded so as to fit into the interior of a funnel (Fig. 1). The paper retains the solid but allows the clear liquid, known as the filtrate, to drain through. The iron particles left upon the filter paper will be found to be identical with[Pg 7] the original iron. The salt can be recovered from the filtrate by evaporation of the water. To accomplish this the filtrate is poured into a small evaporating dish and gently heated (Fig. 2) until the water has disappeared, or evaporated. The solid left in the dish is identical in every way with the original salt. Both the iron and the salt have thus been recovered in their original condition. It is evident that no new substance has been formed by rubbing the salt and iron together. The product is called a mechanical mixture. Such mixtures are very common in nature, almost all minerals, sands, and soils being examples of this class of substances. It is at once apparent that there is no law regulating the composition of a mechanical mixture, and no two mixtures are likely to have exactly the same composition. The ingredients of a mechanical mixture can usually be separated by mechanical means, such as sifting, sorting, magnetic attraction, or by dissolving one constituent and leaving the other unchanged.

Fig. 2

DEFINITION: A mechanical mixture is one in which the constituents retain their original properties, no chemical action having taken place when they were brought together.

Chemical compounds. If iron filings and powdered sulphur are thoroughly ground together in a mortar, a yellowish-green substance results. It might easily be taken to be a new body; but as in the case of the iron and salt, the ingredients can readily be separated. A magnet draws out the iron. Water does not dissolve the sulphur, but other liquids do, as, for example, the liquid called carbon disulphide.[Pg 8] When the mixture is treated with carbon disulphide the iron is left unchanged, and the sulphur can be obtained again, after filtering off the iron, by evaporating the liquid. The substance is, therefore, a mechanical mixture.

If now a new portion of the mixture is placed in a dry test tube and carefully heated in the flame of a Bunsen burner, as shown in Fig. 3, a striking change takes place. The mixture begins to glow at some point, the glow rapidly extending throughout the whole mass. If the test tube is 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 dissolve sulphur from it. It is a new substance with new properties, resulting from the chemical union of iron and sulphur, and is called iron sulphide. Such substances are called chemical compounds, and differ from mechanical mixtures in that the substances producing them lose their own characteristic properties. We shall see later that the two also differ in that the composition of a chemical compound never varies.

Fig. 3

DEFINITION: A chemical compound is a substance the constituents of which have lost their own characteristic properties, and which cannot be separated save by a chemical change.

Elements. It has been seen that iron sulphide is composed of two entirely different substances,—iron and sulphur. The question arises, Do these substances in turn contain other substances, that is, are they also chemical compounds?[Pg 9] Chemists have tried in a great many ways to decompose them, but all their efforts have failed. Substances which have resisted all efforts to decompose them into other substances are called elements. It is not always easy to prove that a given substance is really an element. Some way as yet untried may be successful in decomposing it into other simpler forms of matter, and the supposed element will then prove to be a compound. Water, lime, and many other familiar compounds were at one time thought to be elements.

DEFINITION: An element is a substance which cannot be separated into simpler substances by any known means.

Kinds of matter. While matter has been grouped in three classes for the purpose of study, it will be apparent that there are really but two distinct kinds of matter, namely, compounds and elements. A mechanical mixture is not a third distinct kind of matter, but is made up of varying quantities of either compounds or elements or both.

Alchemy. In olden times it was thought that some way could be found to change one element into another, and a great many efforts were made to accomplish this transformation. Most of these efforts were directed toward changing the commoner metals into gold, and many fanciful ways for doing this were described. The chemists of that time were called alchemists, and the art which they practiced was called alchemy. The alchemists gradually became convinced that the only way common metals could be changed into gold was by the wonderful power of a magic substance which they called the philosopher's stone, which would accomplish this transformation by its mere touch and would in addition give perpetual youth to its fortunate possessor. No one has ever found such a stone, and no one has succeeded in changing one metal into another.

Number of elements. The number of substances now considered to be elements is not large—about eighty in all. Many of these are rare, and very few of them make any[Pg 10] large fraction of the materials in the earth's crust. Clarke gives the following estimate of the composition of the earth's crust:

A complete list of the elements is given in the Appendix. In this list the more common of the elements are marked with an asterisk. It is not necessary to study more than a third of the total number of elements to gain a very good knowledge of chemistry.

Physical state of the elements. About ten of the elements are gases at ordinary temperatures. Two—mercury and bromine—are liquids. The others are all solids, though their melting points vary through wide limits, from cæsium which melts at 26° to elements which do not melt save in the intense heat of the electric furnace.

Occurrence of the elements. Comparatively few of the elements occur as uncombined substances in nature, most of them being found in the form of chemical compounds. When an element does occur by itself, as is the case with gold, we say that it occurs in the free state or native; when it is combined with other substances in the form of compounds, we say that it occurs in the combined state, or in combination. In the latter case there is usually little about the compound to suggest that the element is present in it; for we have seen that elements lose their own peculiar properties when they enter into combination with other elements. It would never be suspected, for example, that the reddish, earthy-looking iron ore contains iron.[Pg 11]

Names of elements. The names given to the elements have been selected in a great many different ways. (1) Some names are very old and their original meaning is obscure. Such names are iron, gold, and copper. (2) Many names indicate some striking physical property of the element. The name bromine, for example, is derived from a Greek word meaning a stench, referring to the extremely unpleasant odor of the substance. The name iodine comes from a word meaning violet, alluding to the beautiful color of iodine vapor. (3) Some names indicate prominent chemical properties of the elements. Thus, nitrogen means the producer of niter, nitrogen being a constituent of niter or saltpeter. Hydrogen means water former, signifying its presence in water. Argon means lazy or inert, the element being so named because of its inactivity. (4) Other elements are named from countries or localities, as germanium and scandium.

Symbols. In indicating the elements found in compounds it is inconvenient to use such long names, and hence chemists have adopted a system of abbreviations. These abbreviations are known as symbols, each element having a distinctive symbol. (1) Sometimes the initial letter of the name will suffice to indicate the element. Thus I stands for iodine, C for carbon. (2) Usually it is necessary to add some other characteristic letter to the symbol, since several names may begin with the same letter. Thus C stands for carbon, Cl for chlorine, Cd for cadmium, Ce for cerium, Cb for columbium. (3) Sometimes the symbol is an abbreviation of the old Latin name. In this way Fe (ferrum) indicates iron, Cu (cuprum), copper, Au (aurum), gold. The symbols are included in the list of elements given in the Appendix. They will become familiar through constant use.[Pg 12]

Chemical affinity the cause of chemical combination. The agency which causes substances to combine and which holds them together when combined is called chemical affinity. The experiments described in this chapter, however, show that heat is often necessary to bring about chemical action. The distinction between the cause producing chemical action and the circumstances favoring it must be clearly made. Chemical affinity is always the cause of chemical union. Many agencies may make it possible for chemical affinity to act by overcoming circumstances which stand in its way. Among these agencies are heat, light, and electricity. As a rule, solution also promotes action between two substances. Sometimes these agencies may overcome chemical attraction and so occasion the decomposition of a compound.

EXERCISES

1. To what class of changes do the following belong? (a) The melting of ice; (b) the souring of milk; (c) the burning of a candle; (d) the explosion of gunpowder; (e) the corrosion of metals. What test question must be applied in each of the above cases?

2. Give two additional examples (a) of chemical changes; (b) of physical changes.

3. Is a chemical change always accompanied by a physical change? Is a physical 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 evolve heat? (b) When the resulting ice melts, is the total heat change the same or different from that of freezing?

6. Give three examples of each of the following: (a) mechanical mixtures; (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 the agency of heat; of light; of electricity.

[Pg 13]

CHAPTER II

OXYGEN

History. The discovery of oxygen is generally attributed to the English chemist Priestley, who in 1774 obtained the element by heating a compound of mercury and oxygen, known as red oxide of mercury. It is probable, however, that the Swedish chemist Scheele had previously obtained it, although an account of his experiments was not published until 1777. The name oxygen signifies acid former. It was given to the element by the French chemist Lavoisier, since he believed that all acids 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. It occurs both in the free and in the combined state. In the free state it occurs in the air, 100 volumes of dry air containing about 21 volumes of oxygen. In the combined state it forms eight ninths of water and nearly one half of the rocks composing the earth's crust. It is also an important constituent of the compounds which compose plant and animal tissues; 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 mixed is such a difficult matter that in the laboratory it has been found more convenient to prepare it from its compounds. The most important of the laboratory methods are the following:[Pg 14]

1. Preparation from water. Water is a compound, consisting of 11.18% hydrogen and 88.82% oxygen. It is easily separated into these constituents by passing an electric current through it under suitable conditions. The process will be described in the chapter on water. While this method of preparation is a simple one, it is not economical.

2. Preparation from mercuric oxide. This method is of interest, since it is the one which led to the discovery of oxygen. The oxide, which consists of 7.4% oxygen and 92.6% mercury, is placed in a small, glass test tube and heated. The compound is in this way decomposed into mercury which collects on the sides of the glass tube, forming a silvery mirror, and oxygen which, being a gas, escapes from the tube. The presence of the oxygen is shown by lighting the end of a splint, extinguishing the flame and bringing the glowing coal into the mouth of the tube. The oxygen causes the glowing coal to burst into a flame.

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 of changes in which all the oxygen is finally set free, leaving a compound of potassium and chlorine called potassium chloride. The change may be represented as follows:

  / potassium \  (potassium       /potassium \ (potassium
 { chlorine   }   chlorate) =    {            } chloride) + oxygen
  \ oxygen    /                   \ chlorine /

JOSEPH PRIESTLEY (English) (1733-1804)

School-teacher, theologian, philosopher, scientist; friend of Benjamin Franklin; discoverer of oxygen; defender of the phlogiston theory; the first to use mercury in a pneumatic trough, by which means he first isolated in gaseous form hydrochloric acid, sulphur dioxide, and ammonia

[Pg 15]

The evolution of the oxygen begins at about 400°. It has been found, however, that if the potassium chlorate is mixed with about one fourth its weight of manganese dioxide, the oxygen is given off at a much lower temperature. Just how the manganese dioxide brings about this result is not definitely known. The amount of oxygen obtained from a given weight of potassium chlorate is exactly the same whether the manganese dioxide is present or not. So far as can be detected the manganese dioxide undergoes no change.

Fig. 4

Directions for preparing oxygen. The manner of preparing oxygen from potassium chlorate is illustrated in the accompanying diagram (Fig. 4). A mixture consisting of one part of manganese dioxide and four parts of potassium chlorate is placed in the flask A and gently heated. The oxygen is evolved and escapes through the tube B. It is collected by bringing over the end of the tube the mouth of a bottle completely filled with water and inverted in a vessel of water, as shown in the figure. The gas rises in the bottle and displaces the water. In the preparation of large quantities of oxygen, a copper retort (Fig. 5) is often substituted for the glass flask.

Fig. 5

In the preparation of oxygen from potassium chlorate and manganese dioxide, the materials used must be pure, otherwise a violent explosion may occur. The purity of the materials is tested by heating a small amount of the mixture in a test tube.

The collection of gases. The method used for collecting oxygen illustrates the general method used for collecting such gases as are[Pg 16] insoluble in water or nearly so. The vessel C (Fig. 4), containing the water in which the bottles are inverted, is called a pneumatic trough.

Commercial methods of preparation. Oxygen can now be purchased stored under great pressure in strong steel cylinders (Fig. 6). It is prepared either by heating a mixture of potassium chlorate and manganese dioxide, or by separating it from the nitrogen and other gases with which it is mixed in the atmosphere. The methods employed for effecting this separation will be described in subsequent chapters.

Fig. 6

Physical properties. Oxygen is a colorless, odorless, tasteless gas, slightly heavier than air. One liter of it, measured at a temperature of 0° and under a pressure of one atmosphere, weighs 1.4285 g., while under similar conditions one liter of air weighs 1.2923 g. It is but slightly soluble in water. Oxygen, like other gases, may be liquefied by applying very great pressure to the highly cooled gas. When the pressure is removed the liquid oxygen passes again into the gaseous state, since its boiling point under ordinary atmospheric pressure is -182.5°.

Chemical properties. At ordinary temperatures oxygen is not very active chemically. Most substances are either not at all affected by it, or the action is so slow as to escape notice. At higher temperatures, however, it is very active, and unites directly with most of the elements. This activity may be shown by heating various substances until just ignited and then bringing them into vessels of the gas, when they will burn with great brilliancy. Thus a glowing splint introduced into a jar of oxygen bursts into flame. Sulphur burns in the air with a very weak flame and feeble light; in oxygen, however, the flame is increased in size and[Pg 17] brightness. Substances which readily burn in air, such as phosphorus, burn in oxygen with dazzling brilliancy. Even substances which burn in air with great difficulty, such as iron, readily burn in oxygen.

The burning of a substance in oxygen is due to the rapid combination of the substance or of the elements composing it with the oxygen. Thus, when sulphur burns both the oxygen and sulphur disappear as such and there is formed a compound of the two, which is an invisible gas, having the characteristic odor of burning sulphur. Similarly, phosphorus on burning forms a white solid compound of phosphorus and oxygen, while iron forms a reddish-black compound of iron and oxygen.

Oxidation. The term oxidation is applied to the chemical change which takes place when a substance, or one of its constituent parts, combines with oxygen. This process may take place rapidly, as in the burning of phosphorus, or slowly, as in the oxidation (or rusting) of iron when exposed to the air. It is always accompanied by the liberation of heat. The amount of heat liberated by the oxidation of a definite weight of any given substance is always the same, being entirely independent of the rapidity of the process. If the oxidation takes place slowly, the heat is generated so slowly that it is difficult to detect it. If the oxidation takes place rapidly, however, the heat is generated in such a short interval of time that the substance may become white hot or burst into a flame.

Combustion; kindling temperature. When oxidation takes place so rapidly that the heat generated is sufficient to cause the substance to glow or burst into a flame the process is called combustion. In order that any substance may undergo combustion, it is necessary that it should be[Pg 18] heated to a certain temperature, known as the kindling temperature. This temperature varies widely for different bodies, but is always definite for the same body. Thus the kindling temperature of phosphorus is far lower than that of iron, but is definite for each. When any portion of a substance is heated until it begins to burn the combustion will continue without the further application of heat, provided the heat generated by the process is sufficient to bring other parts of the substance to the kindling temperature. On the other hand, if the heat generated 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, the compounds formed are called respectively oxide of sulphur, oxide of phosphorus, and oxide of iron. In general, then, an oxide is a compound of oxygen with another element. A great many substances of this class are known; in fact, the oxides of all the common elements have been prepared, with the exception of those of fluorine and bromine. Some of these are familiar compounds. Water, for example, is an oxide of hydrogen, and lime an oxide of the metal calcium.

Products of combustion. The particular oxides formed by the combustion of any substance are called products of combustion of that substance. Thus oxide of sulphur is the product of the combustion of sulphur; oxide of iron is the product of the combustion of iron. It is evident that the products of the combustion of any substance must weigh more than the original substance, the increase in weight corresponding to the amount of oxygen taken up in the act of combustion. For example, when iron burns the oxide of iron formed weighs more than the original iron.[Pg 19]

In some cases the products of combustion are invisible gases, so that the substance undergoing combustion is apparently destroyed. Thus, when a candle burns it is consumed, and so far as the eye can judge nothing is formed during combustion. That invisible gases are formed, however, and that the weight of these is greater than the weight of the candle may be shown by the following experiment.

Fig. 7

Combustion in the broad sense. According to the definition given above, the presence of oxygen is necessary for combustion. The term is sometimes used, however, in a broader sense to designate any chemical change attended by the evolution of heat and light. Thus iron and sulphur, or hydrogen and chlorine under certain conditions, will combine so rapidly that light is evolved, and the action is called a combustion. Whenever combustion takes place in the air, however, the process is one of oxidation.

[Pg 21]

Importance of oxygen. 1. Oxygen is essential to life. Among living organisms only certain minute forms of plant life can exist without it. In the process of respiration the air is taken into the lungs where a certain amount of oxygen is absorbed by the blood. It is then carried to all parts of the body, oxidizing the worn-out tissues and changing them into substances which may readily be eliminated from the body. The heat generated by this oxidation is the source of the heat of the body. The small amount of oxygen which water dissolves from the air supports all the varied forms of aquatic animals.

2. Oxygen is also essential to decay. The process of decay is really a kind of oxidation, but it will only take place in the presence of certain minute forms of life known as bacteria. Just how these assist in the oxidation is not known. By this process the dead products of animal and vegetable life which collect on the surface of the earth are slowly oxidized and so converted into harmless substances. In this way oxygen acts as a great purifying agent.

3. Oxygen is also used in the treatment of certain diseases in which the patient is unable to inhale sufficient air to supply the necessary amount of oxygen.

OZONE

Preparation. When electric sparks are passed through oxygen or air a small percentage of the oxygen is converted into a substance called ozone, which differs greatly from oxygen in its properties. The same change can also be brought about by certain chemical processes. Thus, if some pieces of phosphorus are placed in a bottle and partially covered with water, the presence of ozone may soon be detected in the air contained in the bottle. The conversion of oxygen into ozone is attended by a change in volume, 3 volumes of oxygen forming 2 volumes of ozone. If the resulting ozone is heated to about 300°, the[Pg 22] reverse change takes place, the 2 volumes of ozone being changed back into 3 volumes of oxygen. It is possible that traces of ozone exist in the atmosphere, although its presence there has not been definitely proved, the tests formerly used for its detection having been shown to be unreliable.

Properties. As commonly prepared, ozone is mixed with a large excess of oxygen. It is possible, however, to separate the ozone and thus obtain it in pure form. The gas so obtained has the characteristic odor noticed about electrical machines when in operation. By subjecting it to great pressure 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 chemical properties are similar to those of oxygen except that it is far more active. Air or oxygen containing a small amount of ozone is now used in place of oxygen in certain manufacturing processes.

The difference between oxygen and ozone. Experiments show that in changing oxygen into ozone no other kind of matter is either added to the oxygen or withdrawn from it. The question arises then, How can we account for the difference in their properties? It must be remembered that 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 its properties. That oxygen and ozone contain different amounts of energy may be shown in a number of ways; for example, by the fact that the conversion of ozone into oxygen is attended by the liberation of heat. The passage of the electric sparks through oxygen has in some way changed the energy content of the element and thus it has acquired new properties. Oxygen and ozone must, therefore, be regarded as identical so 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 may exist in more than one form. These different forms of the same element are called allotropic states or forms of the element. These forms differ not only in physical properties but also in their energy contents. Elements often exist in a variety of forms which look quite different. These differences may be due to accidental causes, such as the size or shape of the particles or the way in which the element was prepared. Only such forms, however, as have different energy contents are properly called allotropic forms.[Pg 23]

MEASUREMENT OF GAS VOLUMES

Standard conditions. It is a well-known fact that the volume occupied by a definite weight of any gas can be altered by changing the temperature of the gas or the pressure to which it is subjected. In measuring the volume of gases it is therefore necessary, for the sake of accuracy, to adopt some standard conditions of temperature and pressure. The conditions agreed upon are (1) a temperature of 0°, and (2) a pressure equal to the average pressure exerted by the atmosphere at the sea level, that is, 1033.3 g. per square centimeter. These conditions of temperature and pressure are known as the standard conditions, and when the volume of a gas is given it is understood that the measurement was made under these conditions, unless it is expressly stated otherwise. For example, the weight of a liter of oxygen has been given as 1.4285 g. This means that one liter of oxygen, measured at a temperature of 0° and under a pressure of 1033.3 g. per square centimeter, weighs 1.4285 g.

The conditions which prevail in the laboratory are never the standard conditions. It becomes necessary, therefore, to find a way to calculate the volume which a gas will occupy under standard conditions from the volume which it occupies under any other conditions. This may be done in accordance with the following laws.

Law of Charles. This law expresses the effect which a change in the temperature 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 volume of the gas is increased by 1/273 of the volume which it occupies at zero; likewise for every degree the temperature of the gas falls below zero the volume of the gas is decreased by 1/273 of the volume which it occupies at zero, provided in both cases that the pressure to which the gas is subjected remains constant.

If V represents the volume of gas at 0°, then the volume at 1° will be V + 1/273 V; at 2° it will be V + 2/273 V; or, in general, the volume v, at the temperature t, will be expressed by the formula

Since 1/273 = 0.00366, the formula may be written

Since the value of V (volume under standard conditions) is the one usually sought, it is convenient to transpose the equation to the following form:

The following problem will serve as an illustration of the application of this equation.

The volume of a gas at 20° is 750 cc.; find the volume it will occupy at 0°, the pressure remaining constant.

In this case, v = 750 cc. and t = 20. By substituting these values, equation (4) becomes

Law of Boyle. This law expresses the relation between the volume occupied by a gas and the pressure to which it is subjected. It may be stated as follows: The volume of a gas is inversely proportional to the pressure under which it is measured, provided the temperature of the gas remains constant.

If V represents the volume when subjected to a pressure P and v represents its volume when the pressure is changed to p, then, in accordance 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 numbers representing its volume and the pressure to which it is subjected is a constant.

Since the pressure of the atmosphere at any point is indicated by the barometric reading, it is convenient in the solution of the problems to substitute the latter for the pressure measured in grams per square centimeter. The average reading of the barometer at the sea level is 760 mm., which corresponds to a pressure of 1033.3 g. per square centimeter. The following problem will serve as an illustration of the application of Boyle's law.

A gas occupies a volume of 500 cc. in a laboratory where the barometric reading is 740 mm. What volume would it occupy if the atmospheric pressure 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 and pressure. Inasmuch as corrections must be made as a rule[Pg 25] for both temperature and pressure, it is convenient to combine the equations given above for the corrections for each, so that the two corrections may be made in one operation. The following equation is thus obtained:

in which V_s represents the volume of a gas under standard conditions and v, p, and t the volume, pressure, and temperature respectively at which the gas was actually measured.

The following problem will serve to illustrate the application of this equation.

A gas having a temperature of 20° occupies a volume of 500 cc. when subjected 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 these values in the above equation, we get

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 the preparation of oxygen. In such cases there is present in the gas a certain amount of water vapor. This vapor exerts a definite pressure, which acts in opposition to the atmospheric pressure and which therefore must be subtracted from the latter in determining the effective pressure upon the gas. Thus, suppose we wish to determine the pressure to which the gas in tube A (Fig. 8) is subjected. The tube is raised or lowered until the level of the water inside and outside the tube is the same. The atmosphere presses down upon the surface of the water (as indicated by the arrows), thus forcing the water upward within the tube with a pressure equal to the atmospheric pressure. The full force of this upward pressure, however, is not spent in compressing the gas within the tube, for since it is collected over water it contains a certain amount of water vapor. This water vapor exerts a pressure (as indicated by the arrow within the tube) in opposition to[Pg 26] the upward pressure. It is plain, therefore, that the effective pressure upon the gas is equal to the atmospheric pressure less the pressure exerted by the aqueous vapor. The pressure exerted by the aqueous vapor increases with the temperature. The figures representing the extent of this pressure (often called the tension of aqueous vapor) are given in the Appendix. They express the pressure or tension in millimeters of mercury, just as the atmospheric pressure is expressed in millimeters of mercury. Representing the pressure of the aqueous vapor by a, formula (5) becomes

The following problem will serve to illustrate the method of applying the correction for the pressure of the aqueous vapor.

The volume of a gas measured over water in a laboratory where the temperature is 20° and the barometric reading is 740 mm. is 500 cc. What volume 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. in height. Substituting the values of v, t, p, and a in formula (6), we have

Adjustment of tubes before reading gas volumes. In measuring the volumes of gases collected in graduated tubes or other receivers, over a liquid as illustrated in Fig. 8, the reading should be taken after raising or lowering the tube containing the gas until the level of the liquid inside and outside the tube is the same; for it is only under these conditions that the upward pressure within the tube is the same as the atmospheric 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?[Pg 27]

5. (a) What weight of oxygen can be obtained from 500g. of mercuric oxide? (b) What volume would this occupy under standard conditions?

6. What weight of each of the following compounds is necessary to prepare 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°. At what temperature will its volume be doubled?

9. Reduce the following volumes to standard conditions of pressure, the temperature 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 temperature and pressure: (a) 340 cc. at 12° and 753 mm; (b) 500 cc. at 15° and 740 mm.

12. What weight of potassium chlorate is necessary to prepare 250 l. of oxygen at 20° and 750 mm.?

13. Assuming the cost of potassium chlorate and mercuric oxide to be respectively $0.50 and $1.50 per kilogram, calculate the cost of materials necessary for the preparation of 50 l. of oxygen from each of the above compounds.

14. 100 g. of potassium chlorate and 25 g. of manganese dioxide were heated in the preparation of oxygen. What products were left in the flask, and how much of each was present?

[Pg 28]

CHAPTER III

HYDROGEN

Historical. The element hydrogen was first clearly recognized as a distinct substance by the English investigator Cavendish, who in 1766 obtained it in a pure state, and showed it to be different from the other inflammable airs or gases which had long been known. Lavoisier gave it the name hydrogen, signifying water former, since it had been found to be a constituent of water.

Occurrence. In the free state hydrogen is found in the atmosphere, but only in traces. In the combined state it is widely distributed, being a constituent of water as well as of all living organisms, and the products derived from them, such as starch and sugar. About 10% of the human body is hydrogen. Combined with carbon, it forms the substances which constitute petroleum and natural gas.

Preparation from water. Hydrogen can be prepared from water by several methods, the most important of which are the following.

1. By the electric current. As has been indicated in the preparation of oxygen, water is easily separated into its constituents, hydrogen and oxygen, by passing an electric current through it under certain conditions.

2. By the action of certain metals. When brought into contact with certain metals under appropriate conditions,[Pg 29] water gives up a portion or the whole of its hydrogen, its place being taken by the metal. In the case 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 take place at once, sufficient heat being generated to melt the sodium, which runs about on the surface of the water. The change which takes place consists in the displacement of one half of the hydrogen of the water by the 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 in the undecomposed water, and may be obtained by evaporating the solution to dryness. The hydrogen is evolved as a gas and may be collected by suitable apparatus.

Other metals, such as magnesium and iron, decompose water rapidly, but only at higher temperatures. When steam is passed over hot iron, for example, the iron combines with the oxygen of the steam, thus displacing the hydrogen. Experiments show that the change may be represented as follows:

        _        _
       | hydrogen |           _       _                 _        _
iron + | hydrogen |(water) = | iron    |(iron oxide) + | hydrogen |
       |_oxygen  _|          |_oxygen _|               |_hydrogen_|

The iron oxide formed is a reddish-black compound, identical with that obtained by the combustion of iron in oxygen.

Fig. 9

Preparation from acids (usual laboratory method). While hydrogen can be prepared from water, either by the action of the electric current or by the action of certain metals, these methods are not economical and are therefore but little used. In the laboratory hydrogen is generally prepared from compounds known as acids, all of which contain hydrogen. When acids are brought in contact with certain metals, the metals dissolve and set free the hydrogen[Pg 31] of the acid. Although this reaction is a quite general one, it has been found most convenient in preparing hydrogen by this method to use either zinc or iron as the metal and either hydrochloric or sulphuric acid as the acid. Hydrochloric acid is a compound consisting of 2.77% hydrogen and 97.23% chlorine, while sulphuric acid consists of 2.05% hydrogen, 32.70% sulphur, and 65.25% oxygen.

The changes which take place in the preparation of hydrogen from zinc and sulphuric acid (diluted with water) may be represented as follows:

        _        _               _       _
       | hydrogen |(sulphuric   | zinc    |(zinc
zinc + | sulphur  | acid)     = | sulphur | sulphate) + hydrogen
       |_oxygen  _|             |_oxygen _|

In other words, the zinc has taken the place of the hydrogen in sulphuric acid. The resulting compound contains zinc, sulphur, and oxygen, and is known as zinc sulphate. This remains dissolved in the water present in the acid. It may be obtained in the form of a white solid by evaporating the liquid left after the metal has passed into solution.

When zinc and hydrochloric acid are used the following changes take place:

        _        _                  _        _
       | hydrogen |(hydrochloric   | zinc     |(zinc
zinc + |_chlorine_| acid)        = |_chlorine_| chloride) + hydrogen

When iron is used the changes which take place are exactly similar to those just given for zinc.

Fig. 10.

Physical properties. Hydrogen is similar to oxygen in that it is a colorless, tasteless, odorless gas. It is characterized by its extreme lightness, being the lightest of all known substances. One liter of the gas weighs only 0.08984 g. On comparing this weight with that of an equal volume of oxygen, viz., 1.4285 g., the latter is found to be 15.88 times as heavy as hydrogen. Similarly, air is found to be 14.38 times as heavy as hydrogen. Soap bubbles blown with hydrogen rapidly rise in the air. On account of its lightness it is possible to pour it upward from one bottle into another. Thus, if the bottle A (Fig. 11) is filled with hydrogen, placed mouth downward by the side of bottle B,[Pg 33] filled with air, and is then gradually inverted under B as indicated in the figure, the hydrogen will flow upward into bottle B, displacing the air. Its presence in bottle B may then be shown by bringing a lighted splint to the mouth of the bottle, when the hydrogen will be ignited by the flame. It is evident, from this experiment, that in order to retain the gas in an open bottle the bottle must be placed mouth downward.

Fig. 11

Hydrogen is far more difficult to liquefy than any other gas, with the exception of helium, a rare element recently found to exist in the atmosphere. The English scientist Dewar, however, in 1898 succeeded not only in obtaining hydrogen in liquid state but also as a solid. Liquid hydrogen is colorless and has a density of only 0.07. Its boiling point under atmospheric pressure is -252°. Under diminished pressure the temperature has been reduced to -262°. The solubility of hydrogen in water is very slight, being still less than that of oxygen.

Pure hydrogen produces no injurious results when inhaled. Of course one could not live in an atmosphere of the gas, since oxygen is essential to respiration.

Chemical properties. At ordinary temperatures hydrogen is not an active element. A mixture of hydrogen and chlorine, however, will combine with explosive violence at ordinary temperature if exposed to the sunlight. The union can be brought about also by heating. The product formed in either case is hydrochloric acid. Under suitable conditions hydrogen combines with nitrogen to form ammonia, and with sulphur to form the foul-smelling gas, hydrogen sulphide. The affinity of hydrogen for oxygen is so great that[Pg 34] a mixture of hydrogen and oxygen or hydrogen and air explodes with great violence when heated to the kindling temperature (about 612°). Nevertheless under proper conditions hydrogen may be made to burn quietly in either oxygen or air. The resulting hydrogen flame is almost colorless and is very hot. The combustion of the hydrogen is, of course, due to its union with oxygen. The product of the combustion is therefore a compound of hydrogen and oxygen. That this compound is water may be shown easily by experiment.

Fig. 12

[Pg 35]

Precautions. In order to avoid danger it is absolutely necessary to prove that the hydrogen is free from air before igniting it. This can be done by testing small amounts of the escaping gas. A convenient and safe method of doing this is to fill a test tube with the gas by inverting it over the jet. The hydrogen, on account of its lightness, collects in the tube, displacing the air. After holding it over the jet for a few moments in order that it may be filled with the gas, the tube is gently brought, mouth downward, to the flame of a burner placed not nearer than an arm's length from the jet. If the hydrogen is mixed with air a slight explosion occurs, but if pure it burns quietly in the tube. The operation is repeated until the gas burns quietly, when the tube is quickly brought back over the jet for an instant, whereby the escaping hydrogen is ignited by the flame in the tube.

. Fig. 13

A mixture of hydrogen and oxygen is explosive. That a mixture of hydrogen and air is explosive may be shown safely as follows: A cork through which passes a short glass tube about 1 cm. in diameter is fitted 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 is closed with a small rubber stopper and the bell jar filled with hydrogen, the gas being collected over water. When entirely filled with the gas the jar is removed from the water and supported by blocks of wood 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 once lighted. As the hydrogen escapes, the air flows in at the bottom of the jar and mixes with the remaining portion of the hydrogen, so that a mixture of the two soon forms, and a loud explosion results. The explosion is not dangerous, since the bottom of the jar is open, thus leaving room for the expansion of the hot gas.

Since air is only one fifth oxygen, the remainder being inert gases, it may readily be inferred that a mixture of hydrogen with pure oxygen would be far more explosive than a mixture of hydrogen with air. Such mixtures should not be made except in small quantities and by experienced workers.[Pg 36]

Hydrogen does not support combustion. While hydrogen is readily combustible, it is not a supporter of combustion. In other words, substances will not burn in it. This may be shown by bringing a lighted candle supported by a stiff wire into a bottle or cylinder of the pure gas, as shown in Fig. 14. The hydrogen is ignited by the flame of the candle and burns at the mouth of the bottle, where it comes in contact with 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 slowly withdrawn, the candle is relighted as it passes through the layer of burning hydrogen.

Fig. 14

Fig. 15

Reduction. On account of its great affinity for oxygen, hydrogen has the power of abstracting it from many of its compounds. Thus, if a stream of hydrogen, dried by passing through the tube B (Fig. 15), filled with[Pg 37] calcium chloride, is conducted through the tube C containing some copper oxide, heated to a moderate temperature, the hydrogen abstracts the oxygen from the copper oxide. The change may be represented as follows:

The water formed collects in the cold portions of the tube C near its end. In this experiment the copper oxide is said to undergo reduction. Reduction may therefore be defined as the process of withdrawing oxygen from a compound.

Relation of reduction to oxidation. At the same time that the copper oxide is reduced it is clear that the hydrogen is oxidized, for it combines with the oxygen given up by the copper oxide. The two processes are therefore very closely related, and it usually happens that when one substance is oxidized some other substance is reduced. That substance which gives up its oxygen is called an oxidizing agent, while the substance which unites with the oxygen is called a reducing agent.

The oxyhydrogen blowpipe. This is a form of apparatus used for burning hydrogen in pure oxygen. As has been previously stated, the flame produced by the combustion of hydrogen in the air is very hot. It is evident that if pure oxygen is substituted for air, the temperature reached will be much higher, since there are no inert gases to absorb the heat. The oxyhydrogen blowpipe, used to effect this combination, consists of a small tube placed within a larger one, as shown in Fig. 16.

Fig. 16

[Pg 38]

The hydrogen, stored under pressure, generally in steel cylinders, is first passed through the outer tube and ignited at the open end of the tube. The oxygen from a similar cylinder is then conducted through the inner tube, and mixes with the hydrogen at the end of the tube. In order to produce the maximum heat, the hydrogen and oxygen must be admitted to the blowpipe in the exact proportion in which they combine, viz., 2 volumes of hydrogen to 1 of oxygen, or by weight, 1 part of hydrogen to 7.94 parts of oxygen. The intensity of the heat may be shown by bringing into the flame pieces of metal such as iron wire or zinc. These burn with 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 (calcium oxide), the heat is so intense that the lime becomes incandescent and glows with a brilliant light. This is sometimes used as a source of light, under the name of Drummond or lime light.

Fig. 17

The blast lamp. A similar form of apparatus is commonly used in the laboratory as a source of heat under the name blast lamp (Fig. 17). This differs from the oxyhydrogen blowpipe only in the size of the tubes. In place of the hydrogen and oxygen the more accessible coal gas and air are respectively used. The former is composed largely of a mixture of free hydrogen and gaseous compounds of carbon and hydrogen. While the temperature of the flame is not so high as that of the oxyhydrogen blowpipe, it nevertheless suffices for most chemical operations carried out in the laboratory.

Uses of hydrogen. On account of its cost, hydrogen is but little used for commercial purposes. It is sometimes used as a material for the inflation of balloons, but usually the much cheaper coal gas is substituted for it. Even hot air is often used when the duration of ascension is very short. It has been used also as a source of heat and light[Pg 39] in the oxyhydrogen blowpipe. Where the electric current is available, however, this form of apparatus has been displaced almost entirely by the electric light and electric furnace, which are much more economical and more powerful sources of light and heat.

EXERCISES

1. Will a definite weight of iron decompose an unlimited weight of steam?

2. Why is oxygen passed through the inner tube of the oxyhydrogen blowpipe 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 the quantity 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 use of lime in this instrument?

8. In Fig. 12, why is it necessary to dry the hydrogen by means of the calcium chloride in the tube X?

9. At what pressure would the weight of 1 l. of hydrogen be equal to that of oxygen under standard conditions?

10. (a) What weight of hydrogen can be obtained from 150 g. of sulphuric acid? (b) What volume would this occupy under standard conditions? (c) The density of sulphuric acid is 1.84. What volume would the 150 g. of the acid occupy?

11. How many liters of hydrogen can be obtained from 50 cc. of sulphuric acid having a density of 1.84?

12. Suppose you wish to fill five liter bottles with hydrogen, the gas to be collected over water in your laboratory, how many cubic centimeters of sulphuric acid would be required?

[Pg 40]

CHAPTER IV

COMPOUNDS OF HYDROGEN AND OXYGEN; WATER AND HYDROGEN DIOXIDE

WATER

Historical. Water was long regarded as an element. In 1781 Cavendish showed that it is formed by the union of hydrogen and oxygen. Being a believer in the phlogiston theory, however, he failed to interpret his results correctly. A few years later Lavoisier repeated Cavendish's experiments and showed that water must be regarded as a compound of hydrogen and oxygen.

General methods employed for the determination of the composition of a compound. The composition of a compound may be determined by either of two general processes these are known as analysis and synthesis.

1. Analysis is the process of decomposing a compound into its constituents and determining what these constituents are. The analysis is qualitative when it results in merely determining what elements compose the compound; it is quantitative when the exact percentage of each constituent is determined. Qualitative analysis must therefore precede quantitative analysis, for it must be known what elements, are in a compound before a method can be devised for determining exactly how much of each is present.

2. Synthesis is the process of forming a compound from its constituent parts. It is therefore the reverse of analysis. Like analysis, it may be either qualitative or quantitative.[Pg 41]

Application of these methods to the determination of the composition of water. The determination of the composition of water is a matter of great interest not only because of the importance of the compound but also because the methods employed illustrate the general methods of analysis and synthesis.

Methods based on analysis. The methods based on analysis may be either qualitative or quantitative in character.

Fig. 18

1. Qualitative analysis. As was stated in the study of oxygen, water may be separated into its component parts by means of the electric current. The form of apparatus ordinarily used for effecting this analysis is shown in Fig. 18. A platinum wire, to the end of which is attached a small piece of platinum foil (about 15 mm. by 25 mm.), is fused through each of the tubes B and D, as shown in the figure. The stopcocks at the ends of these tubes are opened and water, to which has been added about one tenth of its volume of sulphuric acid, is poured into the tube A until the side tubes B and D are completely filled. The stopcocks are then closed. The platinum wires extending into the tubes B and D are now connected with the wires leading from two or three dichromate cells joined in series. The pieces of platinum foil within the tubes thus become the electrodes, and the current flows from one to the other through the acidulated water. As soon as the current passes, bubbles of gas rise from each of the electrodes and collect in the upper part of the tubes. The gas[Pg 42] rising from the negative electrode is found to be hydrogen, while that from the positive electrode is oxygen. It will be seen that the volume of the hydrogen is approximately double that of the oxygen. Oxygen is more soluble in water than hydrogen, and a very little of it is also lost by being converted into ozone and other substances. It has been found that when the necessary corrections are made for the error due to these facts, the volume of the hydrogen is exactly double that of the oxygen.

Fig. 19 illustrates a simpler form of apparatus, which may be used in place of that shown in Fig. 18. A glass or porcelain dish is partially filled with water to which has been added the proper amount of acid. Two tubes filled with the same liquid are inverted over the electrodes. The gases resulting from the decomposition of the water collect in the tubes.

Fig. 19

2. Quantitative analysis. The analysis just described is purely qualitative and simply shows that water contains hydrogen and oxygen. It does not prove the absence of other elements; indeed it does not prove that the hydrogen and oxygen are present in the proportion in which they are liberated by the electric current. The method may be made quantitative, however, by weighing the water decomposed and also the hydrogen and oxygen obtained in its decomposition. If the combined weights of the hydrogen and oxygen exactly equal the weight of the water decomposed, then it would[Pg 43] be proved that the water consists of hydrogen and oxygen in the proportion in which they are liberated by the electric current. This experiment is difficult to carry out, however, so that the more accurate methods based on synthesis are used.

Methods based on synthesis. Two steps are necessary to ascertain the exact composition of water by synthesis: (1) to show by qualitative synthesis that water is formed by the union of oxygen with hydrogen; (2) to determine by quantitative synthesis in what proportion the two elements unite to form water. The fact that water is formed by the combination of oxygen with hydrogen was proved in the preceding chapter. The quantitative synthesis may be made as follows:

Fig. 20

The combination of the two gases is brought about in a tube called a eudiometer. 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 are fused through the glass, the ends of the wires within the tube being separated by a space of 2 mm or 3 mm. The tube is entirely filled with mercury and inverted in a vessel of the same liquid. Pure hydrogen is passed into the tube until it is about one fourth filled. The volume of the gas is then read off on the scale and reduced to standard conditions. Approximately an equal volume of pure oxygen is then introduced and the volume again read off and reduced to standard conditions. This gives the total volume of the two gases. From this the volume of the oxygen introduced may be determined by[Pg 44] subtracting from it the volume of the hydrogen. The combination of the two gases is now brought about by connecting the two platinum wires with an induction coil and passing a spark from one wire to the other. Immediately a slight explosion occurs. The mercury in the tube is at first depressed because of the expansion of the gases due to the heat generated, but at once rebounds, taking the place of the gases which have combined to form water. The volume of the water in the liquid state is so small that it may be disregarded in the calculations. In order that the temperature of the residual gas and the mercury may become uniform, the apparatus is allowed to stand for a few minutes. The volume of the gas is then read off and reduced to standard conditions, so that it may be compared with the volumes of the hydrogen and oxygen originally taken. The residual gas is then tested in order to ascertain whether it is hydrogen or oxygen, experiments having proved that it is never a mixture of the two. From the information thus obtained the composition of the water may be calculated. Thus, suppose the readings were as follows:

The 20.3 cc. of hydrogen have combined with 18.4 cc. minus 8.3 cc. (or 10.1 cc.) of oxygen; or approximately 2 volumes of hydrogen have combined with 1 of oxygen. Since oxygen is 15.88 times as heavy as hydrogen, the proportion by weight in which the two gases combine is 1 part of hydrogen to 7.94 of oxygen.[Pg 45]

Precaution. If the two gases are introduced into the eudiometer in the exact proportions in which they combine, after the combination has taken place the liquid will rise and completely fill the tube. Under these conditions, however, the tube is very likely to be broken by the sudden upward rush of the liquid. Hence in performing the experiment care is taken 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 the calculations necessary in reducing the volumes of the gases to the same conditions of temperature and pressure in order to make comparisons. With this apparatus it is possible to take the readings of the volumes under the same conditions of temperature and pressure, and thus compare them directly. The apparatus (Fig. 21) is filled with mercury and the gases introduced into the tube A. The experiment is carried out as in the preceding one, except that before taking the reading of the gas volumes, mercury is either added to the tube B or withdrawn from it by means of the stopcock C, until it stands at exactly the same height in both tubes. The gas inclosed in tube A is then under atmospheric pressure; and since but a few minutes are required for performing the experiment, the conditions of temperature and pressure may be regarded as constant. Hence the volumes of the hydrogen and oxygen and of the residual gas may be read off from the tube and directly compared.

Fig. 21

Method used by Berzelius and Dumas. The method used by these investigators enables us to determine directly the proportion by weight in which the hydrogen and oxygen combine. Fig. 22 illustrates the apparatus used in making this determination. B is a glass tube containing copper oxide. C and D are glass tubes filled with calcium chloride, a substance which has great affinity for water.[Pg 46] The tubes B and C, including their contents, are carefully weighed, and the apparatus connected as shown in the figure. A slow current of pure hydrogen is then passed through A, and that part of the tube B which contains copper oxide is carefully heated. The hydrogen combines with the oxygen present in the copper oxide to form water, which is absorbed by the calcium chloride in tube C. The calcium chloride in tube D prevents any moisture entering tube C from the air. The operation is continued until an appreciable amount of water has been formed. The tubes B and C are then weighed once more. The loss of weight in the tube B will exactly equal the weight of oxygen taken up from the copper oxide in the formation of the water. The gain in weight in the tube C will exactly equal the weight of the water formed. The difference in these weights will of course equal the weight of the hydrogen present in the water formed.

Fig. 22

Dumas' results. The above method for the determination of the composition of water was first used by Berzelius in 1820. The work was repeated in 1843 by Dumas, the average of whose results is as follows:

[Pg 47]

According to this experiment the ratio of hydrogen to oxygen in water is therefore 26.32 to 210.04, or as l to 7.98

Morley's results. The American chemist Morley has recently determined the composition of water, extreme precautions being taken to use pure materials and to eliminate all sources of error. The hydrogen and oxygen which combined, as well as the water formed, were all accurately weighed. According to Morley's results, 1 part of hydrogen by weight combines with 7.94 parts of oxygen to form water.

Comparison of results obtained. From the above discussions it is easy to see that it is by experiment alone that the composition of a compound can be determined. Different methods may lead to slightly different results. The more accurate the method chosen and the greater the skill with which the experiment is carried out, the more accurate will be the results. It is generally conceded by chemists that the results obtained by Morley in reference to the composition of water are the most accurate ones. In accordance with these results, then, water must be regarded as a compound containing hydrogen and oxygen in the proportion of 1 part by weight of hydrogen to 7.94 parts by weight of oxygen.

Relation between the volume of aqueous vapor and the volumes of the hydrogen and oxygen which combine to form it. When the quantitative synthesis of water is carried out in the eudiometer as described above, the water vapor formed by the union of the hydrogen and oxygen at once condenses. The volume of the resulting liquid is so small that it may be disregarded in making the calculations. If, however, the experiment is carried out at a temperature of 100° or above, the water-vapor formed is not condensed and it thus becomes possible to compare the volume of the[Pg 48] vapor 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 obtained by boiling some liquid which has a boiling point above 100°. In this way it has been proved that 2 volumes of hydrogen and 1 volume of oxygen combine to form exactly 2 volumes of water vapor, the volumes all being measured under the same conditions of temperature and pressure. It will be noted that the relation between these volumes may be expressed by whole numbers. The significance of this very important fact will be discussed in a subsequent chapter.

Fig. 23

Occurrence of water. Water not only covers about three fourths of the surface of the earth, and is present in the atmosphere in the form of moisture, but it is also a common constituent of the soil and rocks and of almost every form of animal and vegetable organism. The human body is nearly 70% water. This is derived not only from the water which we drink but also from the food which we eat, most of which contains a large percentage 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 and oxygen. Such a water never occurs in nature, however, for being a good solvent, it takes up certain substances from the rocks and soil with which it comes in contact. When such waters are evaporated these[Pg 49] substances are deposited in the form of a residue. Even rain water, which is the purest form occurring in nature, contains dust particles and gases dissolved from the atmosphere. The foreign matter in water is of two kinds, namely, mineral, such as common salt and limestone, and organic, that is the products of animal and vegetable life.

Effect of impurities on health. The mineral matter in water does not, save in very exceptional cases, render the water injurious to the human system. In fact the presence of a certain amount of such matter is advantageous, supplying the mineral constituents necessary for the formation of the solid tissues of the body. The presence of organic matter, on the other hand, must always be regarded with suspicion. This organic matter may consist not only of the products of animal and vegetable life but also of certain microscopic forms of living organisms which are likely to accompany such products. Contagious diseases are known to be due to the presence in the body of minute living organisms or germs. Each disease is caused by its own particular kind of germ. Through sewage these germs may find their way from persons afflicted with disease into the water supply, and it is principally through the drinking water that certain of these diseases, especially typhoid fever, are spread. It becomes of great importance, therefore, to be[Pg 50] able to detect such matter when present in drinking water as well as to devise methods whereby it can be removed or at least rendered harmless.

Fig. 24

Purification of water. Three general methods are used for the purification of water, namely, distillation, filtration, and boiling.

1. Distillation. The most effective way of purifying natural waters is by the process of distillation. This consists in boiling the water and condensing the steam. Fig. 24 illustrates the process of distillation, as commonly conducted[Pg 51] in the laboratory. Ordinary water is poured into the flask A and boiled. The steam is conducted through the condenser B, which consists essentially of a narrow glass tube sealed within a larger 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 kept cool and the steam in passing through it is condensed. The water formed by the condensation of the steam collects in the receiver E and is known as distilled water. Such water is practically pure, since the impurities are nonvolatile and remain in the flask A.

2. Filtration. The process of distillation practically removes all nonvolatile foreign matter, mineral as well as organic. In purifying water for drinking purposes, however, it is only necessary to eliminate the latter or to render it harmless. This is ordinarily done either by filtration or[Pg 52] boiling. In filtration the water is passed through some medium which will retain the organic matter. Ordinary charcoal is a porous substance and will condense within its pores the organic matter in water if brought in contact with it. It is therefore well adapted to the construction of filters. Such filters to be effective must be kept clean, since it is evident that the charcoal is useless after its pores are filled. A more effective type of filter is the Chamberlain-Pasteur filter. In this the water is forced through a porous cylindrical cup, the pores being so minute as to strain out the organic matter.

Fig. 25

3. Boiling. A simpler and equally efficient method for purifying water for drinking purposes consists in boiling the water. It is the germs in water that render it dangerous to health. These germs are living forms of matter. If the[Pg 53] water is boiled, the germs are killed and the water rendered safe. While these germs are destroyed by heat, cold has little effect upon them. Thus Dewar, in working with liquid hydrogen, exposed some of these minute forms of life to the temperature of boiling hydrogen (-252°) without killing them.

Self-purification of water. It has long been known that water contaminated with organic matter tends to purify itself when exposed to the air. This is due to the fact that the water takes up a small amount of oxygen from the air, which gradually oxidizes the organic matter present in the water. While water is undoubtedly purified in this way, the method cannot be relied upon to purify a contaminated water so as to render 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 observed through a considerable thickness. It solidifies at 0° and boils at 100° under the normal pressure of one atmosphere. If the pressure is increased, the boiling point is raised. When water is cooled it steadily contracts until the temperature of 4° is reached: it then expands. Water is remarkable for its ability to dissolve other substances, and is the best solvent known. Solutions of solids in water are more frequently employed in chemical work than are the solid substances, for chemical action between substances goes on more readily when they are in solution than it does when they are in the solid state.

Chemical properties. Water is a very stable substance, or, in other words, it does not undergo decomposition readily. To decompose it into its elements by heat alone requires a very high temperature; at 2500°, for example, only about 5% of the entire amount is decomposed. Though very[Pg 54] stable towards heat, water can be decomposed in other ways, as by the action of the electrical current or by certain metals.

Water of crystallization. When a solid is dissolved in water and the resulting solution is allowed to evaporate, the solid separates out, often in the form of crystals. It has been found that the crystals of many compounds, although perfectly dry, give up a definite amount of water when heated, the substance at the same time losing its crystalline form. Such water is called water of crystallization. This varies in amount with different compounds, but is perfectly definite for the same compound. Thus, if a perfectly dry crystal of copper sulphate is strongly heated in a tube, water is evolved and condenses on the sides of the tube, the crystal crumbling to a light powder. The weight of the water evolved is always equal to exactly 36.07% of the weight of copper sulphate crystals heated. The water must therefore be in chemical combination with the substance composing the crystal; for if simply mixed with it or adhering to it, not only would the substance appear moist but the amount present would undoubtedly vary. The combination, however, must be a very weak one, since the water is often expelled by even a gentle heat. Indeed, in some cases the water is given up on simple exposure to air. Such compounds are said to be efflorescent. Thus a crystal of sodium sulphate[Pg 55] (Glauber's salt) on exposure to air crumbles to a fine powder, owing to the escape of its water of crystallization. Other substances have just the opposite property: they absorb moisture when exposed to the air. For example, if a bit of dry calcium chloride is placed in moist air, in the course of a few hours it will have absorbed sufficient moisture to dissolve it. Such substances are said to be deliquescent. A deliquescent body serves as a good drying or desiccating agent. We have already employed calcium chloride as an agent for absorbing the moisture from hydrogen. Many substances, as for example quartz, form crystals which contain no water of crystallization.

Uses of water. The importance of water in its relation to life and commerce is too well known to require comment. Its importance to the chemist has also been pointed out. It remains to call attention to the fact that it is used as a standard in many physical measurements. Thus 0° and 100° on the centigrade scale are respectively the freezing and the 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 the metric system, namely, the gram. It is also taken as the unit for the determination of the density of liquids and solids as well as for the measurement of amounts of heat.[Pg 56]

HYDROGEN DIOXIDE

Composition. As has been shown, 1 part by weight of hydrogen combines with 7.94 parts by weight of oxygen to form water. It is possible, however, to obtain a second compound of hydrogen and oxygen differing from water in composition in that 1 part by weight of hydrogen is combined with 2 × 7.94, or 15.88 parts, of oxygen. This compound is called 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 direct union of hydrogen and oxygen, and indirect methods must therefore be used. It is commonly prepared by the action of a solution of sulphuric acid on barium dioxide. The change which takes place may be indicated as follows:

In other words, the barium and hydrogen in the two compounds exchange places. By this method a dilute solution of the dioxide in water is obtained. It is possible to separate the dioxide from the water by fractional distillation. This is attended with great difficulties, however, since the pure dioxide is explosive. The distillation is carried on under diminished pressure so as to lower the boiling points as much as possible; otherwise the high temperature would decompose the dioxide.[Pg 57]

Properties. Pure hydrogen dioxide is a colorless sirupy liquid having a density of 1.49. Its most characteristic property is the ease with which it decomposes into water and oxygen. One part by weight of hydrogen is capable of holding firmly only 7.94 parts of oxygen. The additional 7.94 parts of oxygen present in hydrogen dioxide are therefore easily evolved, the compound breaking down into water and oxygen. This decomposition is attended by the generation of considerable heat. In dilute solution hydrogen dioxide is fairly stable, although such a solution should be kept in a dark, cool place, since both heat and light aid in the decomposition of the dioxide.

Uses. Solutions of hydrogen dioxide are used largely as oxidizing agents. The solution sold by druggists contains 3% of the dioxide and is used in medicine as an antiseptic. Its use as an antiseptic depends upon its oxidizing properties.

EXERCISES

1. Why does the chemist use distilled water in making solutions, rather than filtered water?

2. How could you determine the total amount of solid matter dissolved in a sample of water?

3. How could you determine whether a given sample of water is distilled water?

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 be determined?

6. Would ice frozen from impure water necessarily be free from disease germs?

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 inclosed water and water of crystallization?[Pg 58]

9. If steam is heated to 2000° and again cooled, has any chemical change taken 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 a glass condenser.

12. Draw a diagram of the apparatus used in your laboratory for supplying distilled water.

13. 20 cc. of hydrogen and 7 cc. of oxygen are placed in a eudiometer and the mixture exploded. (a) How many cubic centimeters of aqueous vapor 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 L of hydrogen, measured under standard conditions? (b)What volume of oxygen would be required in (a)? (c)What weight of potassium chlorate is necessary to prepare this amount of oxygen?

15. What weight of oxygen is present in 1 kg. of the ordinary hydrogen dioxide solution? In the decomposition of this weight of the dioxide into water and oxygen, what volume of oxygen (measured under standard conditions) is evolved?

[Pg 59]

CHAPTER V

THE ATOMIC THEORY

Three fundamental laws of matter. Before we can gain any very definite idea in regard to the structure of matter, and the way in which different kinds of substances act chemically upon each other, it is necessary to have clearly in view three fundamental laws of matter. These laws have been established by experiment, and any conception which may 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 in the introductory chapter, and needs no further discussion. It will be recalled that it may be stated thus: Matter can neither be created nor destroyed, though it can be changed from one form into another.

Law of definite composition. In the earlier days of chemistry there was much discussion as to whether the composition of a given compound is always precisely the same or whether it is subject to some variation. Two Frenchmen, Berthollet and Proust, were the leaders in this discussion, and a great deal of most useful experimenting was done to decide the question. Their experiments, as well as all succeeding ones, have shown that the composition of a pure chemical compound is always exactly the same. Water obtained by melting pure ice, condensing steam, burning hydrogen in oxygen, has always 11.18% hydrogen and 88.82% oxygen in it. Red oxide of mercury, from whatever source it is obtained, contains 92.6%[Pg 60] mercury and 7.4% oxygen. This truth is known as the law of definite composition, and may be stated thus: The composition of a chemical compound never varies.

Law of multiple proportion. It has already been noted, however, that hydrogen and oxygen combine in two different ratios to form water and hydrogen dioxide respectively. It will be observed that this fact does not contradict the law of definite composition, for entirely different substances are formed. These compounds differ from each other in composition, but the composition of each one is always constant. This ability of two elements to unite in more than one ratio is very frequently 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 English school-teacher and philosopher, endeavored to find some rule which holds between the ratios in which two given substances combine. His studies brought to light a very simple relation, which the following examples will make clear. In water the hydrogen and oxygen are combined in the ratio 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 with 15.88 parts by weight of oxygen. The ratio between the amounts of oxygen which combine with the same amount of hydrogen to form water and hydrogen dioxide respectively is therefore 7.94: 15.88, or 1: 2.

JOHN DALTON (English) (1766-1844)

Developed the atomic theory; made many studies on the properties and the composition of gases. His book entitled "A New System of Chemical Philosophy" had a large influence on the development of chemistry

[Pg 61]

Similarly, the element iron combines with oxygen to form two oxides, one of which is black and the other red. By analysis it has been shown that the former contains 1 part by weight of iron combined with 0.286 parts by weight of oxygen, while the latter contains 1 part by weight of iron combined with 0.429 parts by weight of oxygen. Here again we find that the amounts of oxygen which combine with the same fixed amount of iron to 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 has been found to hold true in all cases where more than one compound is, formed from the same elements. Dalton's law of multiple proportion states these facts as follows: When any two elements, A and B, combine to form more than one compound, the amounts of B which unite with any fixed amount of A bear the ratio of small whole numbers to each other.

Hypothesis necessary to explain the laws of matter. These three generalizations are called laws, because they express in concise language truths which are found by careful experiment to hold good in all cases. They do not offer any explanation of the facts, but merely state them. The human mind, however, does not rest content with the mere bare facts, but seeks ever to learn the explanation of the facts. A suggestion which is offered to explain such a set of facts is called an hypothesis. The suggestion which Dalton offered to explain the three laws of matter, called the atomic hypothesis, was prompted by his view of the constitution of matter, and it involves three distinct assumptions in regard to the nature of matter and chemical action. Dalton could not prove these assumptions to be true, but he saw that if they were true the laws of matter become very easy to understand.

Dalton's atomic hypothesis. The three assumptions which Dalton made in regard to the nature of matter, and which together constitute the atomic hypothesis, are these:[Pg 62]

1. All elements are made up of minute, independent particles which Dalton designated as atoms.

2. All atoms of the same element have equal masses; those of different elements have different masses; in any change to which an atom is subjected its mass does not change.

3. When two or more elements unite to form a compound, the action consists in the union of a definite small number of atoms of each element to form a small particle of the compound. The smallest particles of a given compound are therefore exactly alike in the number and kinds of atoms which they contain, and larger masses of the substances are simply aggregations of these least particles.

Molecules and atoms. Dalton applied the name atom not only to the minute particles of the elements but also to the least particles of compounds. Later Avogadro, an Italian scientist, pointed out the fact that the two are different, since the smallest particle of an element is a unit, while that of a compound must have at least two units in it. He suggested the name molecule for the least particle of a compound which can exist, retaining the name atom for the smallest particle of an element. In accordance with this distinction, we may define the atom and the molecule as follows: An atom is the smallest particle of an element which can exist. A molecule is the smallest particle of a compound which can exist. It will be shown in a subsequent chapter that sometimes two or more atoms of the same element unite with each other to form molecules of the element. While the term atom, therefore, is applicable only to elements, the term molecule is applicable both to elements and compounds.[Pg 63]

The atomic hypothesis and the laws of matter. Supposing the atomic hypothesis to be true, let us now see if it is in harmony with the laws of matter.

1. The atomic hypothesis and the law of conservation of matter. It is evident that if the atoms never change their masses in any change which they undergo, the total quantity of matter can never change and the law of 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 the union is between definite numbers of the two kinds of atoms. In the simplest case one atom of the one element combines with one atom of the other. If the sulphur and the iron atoms never change their respective masses when they unite to form a molecule of iron sulphide, all iron sulphide molecules will have equal amounts of iron in them and also of sulphur. Consequently any mass made up of iron sulphide molecules will have the same fraction of iron by weight as do the individual iron sulphide molecules. Iron sulphide, from whatever source, will have the same composition, which is in accordance with the law of definite composition.

3. The atomic hypothesis and the law of multiple proportion. But this simplest case may not always be the only one. Under other conditions one atom of iron might combine with two of sulphur to form a molecule of a second compound. In such a case the one atom of iron would be in combination with twice the mass of sulphur that is in the first compound, since the sulphur atoms all have equal masses. What is true for one molecule will be true for any number of them; consequently when such quantities of these two compounds are selected as are found to contain[Pg 64] the same amount of iron, the one will contain twice as much sulphur as the other.

The combination between the atoms may of course take place in other simple ratios. For example, two atoms of one element might combine with three or with five of the other. In all such cases it is clear that the law of multiple proportion must hold true. For on selecting such numbers of the two kinds of molecules as have the same number of the one kind of atoms, the numbers of the other kind of atoms will stand in some simple ratio to each other, and their weights will therefore stand in the same simple ratio.

Testing the hypothesis. Efforts have been made to find compounds which do not conform to these laws, but all such attempts have resulted in failure. If such compounds should be found, the laws would be no longer true, and the hypothesis of Dalton would cease to possess value. When an hypothesis 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 is termed a theory. We now speak of the atomic theory rather than of the atomic hypothesis.

Value of a theory. The value of a theory is twofold. It aids in the clear understanding of the laws of nature because it gives an intelligent idea as to why these laws should be in operation.

A theory also leads to discoveries. It usually happens that in testing a theory much valuable work is done, and many new facts are discovered. Almost any theory in explaining given laws will involve a number of consequences apart from the laws it seeks to explain. Experiment will soon show whether these facts are as the theory predicts they will be. Thus Dalton's atomic theory predicted many properties of gases which experiment has since verified.[Pg 65]

Atomic weights. It would be of great advantage in the study of chemistry if we could determine the weights of the different kinds of atoms. It is evident that this cannot be done directly. They are so small that they cannot be seen even with a most powerful microscope. It is calculated that it would take 200,000,000 hydrogen atoms placed side by side to make a row one centimeter long. No balance can weigh such minute objects. It is possible, however, to determine their relative weights,—that is, how much heavier one is than another. These relative weights of the atoms are spoken of as the atomic weights of the elements.

If elements were able to combine in only one way,—one atom of one with one atom of another,—the problem of determining the atomic weights would be very simple. We should merely have to take some one convenient element as a standard, and find by experiment how much of each other element would combine with a fixed weight of it. The ratios thus found would be the same ratios as those between the atoms of the elements, and thus we should have their relative atomic weights. The law of multiple proportion calls attention to the fact that the atoms combine in other ratios than 1: 1, and there is no direct way of telling which one, if any, of the several compounds in a given case is the one consisting of a single atom of each element.

If some way were to be found of telling how much heavier the entire molecule of a compound is than the atom chosen as a standard,—that is, of determining the molecular weights of compounds,—the problem could be solved, though its solution would not be an entirely simple matter. There are ways of determining the molecular weights of[Pg 66] compounds, and there are other experiments which throw light directly upon the relative weights of the atoms. These methods cannot be described until the facts upon which they rest have been studied. It will be sufficient for the present to assume that these methods are trustworthy.

Standard for atomic weights. Since the atomic weights are merely relative to some one element chosen as a standard, it is evident that any one of the elements may serve as this standard and that any convenient value may be assigned to its atom. At one time oxygen was taken as this standard, with the value 100, and the atomic weights of the other elements were expressed in terms of this standard. It would seem more rational to take the element of smallest atomic weight as the standard and give it unit value; accordingly hydrogen was taken as the standard with an atomic weight of 1. Very recently, however, this unit has been replaced by oxygen, with an atomic weight of 16.

Why oxygen is chosen as the standard for atomic weights. In the determination of the atomic weight of an element it is necessary to find the weight of the element which combines with a definite weight of another element, preferably the element chosen as the standard. Since oxygen combines with the elements far more readily than does hydrogen to form definite compounds, it is far better adapted for the standard element, and has accordingly replaced hydrogen as the standard. Any definite value might be given to the weight of the oxygen atom. In assigning a value to it, however, it is convenient to choose a whole number, and as small a number as possible without making the atomic weight of any other element less than unity. For these reasons the number 16 has been chosen as the atomic[Pg 67] weight of oxygen. This makes the atomic weight of hydrogen equal to 1.008, so that there is but little difference between taking oxygen as 16 and hydrogen as 1 for the unit.

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 the same weight?

[Pg 68]

CHAPTER VI

CHEMICAL EQUATIONS AND CALCULATIONS

Formulas. Since the molecule of any chemical compound consists of a definite number of atoms, and this number never changes without destroying the identity of the compound, it is very convenient to represent the composition of a compound by indicating the composition of its molecules. This can be done very easily by using the symbols of the atoms to indicate the number and the kind of the atoms which constitute the molecule. HgO will in this way represent mercuric oxide, a molecule of which has been found to contain 1 atom each of mercury and oxygen. H₂O will represent water, the molecules of which consist of 1 atom of oxygen and 2 of hydrogen, the subscript figure indicating the number of the atoms of the element whose symbol precedes it. H₂SO₄ will stand for sulphuric acid, the molecules of which contain 2 atoms of hydrogen, 1 of sulphur, and 4 of oxygen. The combination of symbols which represents the molecule of a substance is called its formula.

Equations. When a given substance undergoes a chemical change it is possible to represent this change by the use of such symbols and formulas. In a former chapter it was shown that mercuric oxide decomposes when heated to form mercury and oxygen. This may be expressed very briefly in the form of the equation

[Pg 69]

When water is electrolyzed two new substances, hydrogen and oxygen, are formed from it. This statement in the form of an equation is

The coefficient before the symbol for hydrogen indicates that a single molecule of water yields two atoms of hydrogen on decomposition.

In like manner the combination of sulphur with iron is expressed by the equation

The decomposition of potassium chlorate by heat takes place as represented by the equation

Reading of equations. Since equations are simply a kind of shorthand way of indicating chemical changes which occur under certain conditions, in reading an equation the full statement for which it stands should be given. Equation (1) should be read, "Mercuric oxide when heated gives mercury and oxygen"; equation (2) is equivalent to the statement, "When electrolyzed, water produces hydrogen and oxygen"; equation (3), "When heated together iron and sulphur unite to form iron sulphide"; equation (4), "Potassium chlorate when heated yields potassium chloride and oxygen."

Knowledge required for writing equations. In order to write such equations correctly, a considerable amount of exact knowledge is required. Thus, in equation (1) the fact that red oxide of mercury has the composition represented by the formula HgO, that it is decomposed by heat, that in this decomposition mercury and oxygen are formed and[Pg 70] no other products,—all these facts must be ascertained by exact experiment before the equation can be written. An equation expressing these facts will then have much value.

Having obtained an equation describing the conduct of mercuric oxide on being heated, it will not do to assume that other oxides will behave in like manner. Iron oxide (FeO) resembles mercuric oxide in many respects, but it undergoes no change at all when heated. Manganese dioxide, the black substance used in the preparation of oxygen, has the formula MnO₂. When this substance is heated oxygen is set free, but the metal manganese is not liberated; instead, a different oxide of manganese containing less oxygen is produced. The equation representing the reaction is

Classes of reactions. When a chemical change takes place in a substance the substance is said to undergo a reaction. Although a great many different reactions will be met in the study of chemistry, they may all be grouped under the following heads.

1. Addition. This is the simplest kind of chemical action. It consists in the union of two or more substances to produce a new substance. The combination of iron with sulphur is an example:

2. Decomposition. This is the reverse of addition, the substance undergoing reaction being parted into its constituents. The decomposition of mercuric oxide is an example: HgO = Hg + O.

3. Substitution. It is sometimes possible for an element in the free state to act upon a compound in such a way that[Pg 71] it takes the place of one of the elements of the compound, liberating it in turn. In the study of the element hydrogen it was pointed out that hydrogen is most conveniently prepared by the action of sulphuric or hydrochloric acid upon zinc. When sulphuric acid is used a substance called zinc sulphate, having the composition represented by the formula ZnSO₄, is formed together with hydrogen. The equation is

When hydrochloric acid is used zinc chloride and hydrogen are the products of reaction:

When iron is used in place of zinc the equation is

These reactions are quite similar, as is apparent from an examination of the equations. In each case 1 atom of the metal replaces 2 atoms of hydrogen in the acid, and the hydrogen escapes as a gas. When an element in the free state, such as the zinc in the equations just given, takes the place of some one element in a compound, setting it free from chemical combination, the act is called substitution.

Other reactions illustrating substitution are the action of sodium on water,

and the action of heated iron upon water,

4. Double decomposition. When barium dioxide (BaO₂) is treated with sulphuric acid two compounds are formed, namely, hydrogen dioxide (H₂O₂) and barium sulphate (BaSO₄). The equation is

[Pg 72]

In this reaction it will be seen that the two elements barium and hydrogen simply exchange places. Such a reaction is called a double decomposition. We shall meet with many examples of this kind of chemical reactions.

Chemical equations are quantitative. The use of symbols and formulas in expressing chemical changes has another great advantage. Thus, according to the equation

1 molecule of water is decomposed into 2 atoms of hydrogen and 1 atom of oxygen. But, as we have seen, the relative weights of the atoms are known, that of hydrogen being 1.008, while that of oxygen is 16. The molecule of water, being composed of 2 atoms of hydrogen and 1 atom of oxygen, must therefore weigh relatively 2.016 + 16, or 18.016. The amount of hydrogen in this molecule must be 2.016/18.016, or 11.18% of the whole, while the amount of oxygen must be 16/18.018, or 88.82% of the whole. Now, since any definite quantity of water is simply the sum of a great many molecules of water, it is plain that the fractions representing the relative amounts of hydrogen and oxygen present in a molecule must likewise express the relative amounts of hydrogen and oxygen present in any quantity of water. Thus, for example, in 20 g. of water there are 2.016/18.016 × 20, or 2.238 g. of hydrogen, and 16/18.016 × 20, or 17.762 g. of oxygen. These results in reference to the composition of water of course agree exactly with the facts obtained by the experiments described in the chapter on water, for it is because of those experiments that the values 1.008 and 16 are given to hydrogen and oxygen respectively.

It is often easier to make calculations of this kind in the form of a proportion rather than by fractions. Since the[Pg 73] molecule of water and the two atoms of hydrogen which it contains have the ratio by weight of 18.016: 2.016, any mass of water has the same ratio between its total weight and the weight of the hydrogen in it. Hence, to find the number of grams (x) of hydrogen in 20 g. of water, we have the proportion

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. of water, we have the proportion

from which we find that x = 17.762 g.

Again, suppose we wish to find what weight of oxygen can be obtained from 15 g. of mercuric oxide. The equation representing the decomposition of mercuric oxide is

The relative weights of the mercury and oxygen atoms are respectively 200 and 16. The relative weight of the mercuric oxide molecule must therefore be the sum of these, or 216. The molecule of mercuric oxide and the atom of oxygen which it contains have the ratio 216: 16. This same ratio must therefore hold between the weight of any given quantity of mercuric oxide and that of the oxygen which it contains. Hence, to find the weight of oxygen in 15 g. of mercuric oxide, we have the proportion

On the other hand, suppose we wish to prepare, say, 20 g. of oxygen. The problem is to find out what weight of mercuric oxide will yield 20 g. of oxygen. The following proportion evidently holds

from which we get x = 270.

In the preparation of hydrogen by the action of sulphuric acid upon zinc, according to the equation,

[Pg 74]

suppose that 50 g. of zinc are available; let it be required to calculate the weight of hydrogen which can be obtained. It will be seen that 1 atom of zinc will liberate 2 atoms of hydrogen. The ratio by weight of a zinc to an hydrogen atom is 65.4: 1.008; of 1 zinc atom to 2 hydrogen atoms, 65.4: 2.016. Zinc and hydrogen will be related in this reaction in this same ratio, however many atoms of zinc are concerned. Consequently in the proportion

x will be the weight of hydrogen set free by 50 g. of zinc. The weight of zinc sulphate produced at the same time can be found from the proportion

where 161.46 is the molecular weight of the zinc sulphate, and x the weight of zinc sulphate formed. In like manner, the weight of sulphuric acid used up can be calculated from the proportion

These simple calculations are possible because the symbols and formulas in the equations represent the relative weights of the substances concerned in a chemical reaction. When once the relative weights of the atoms have been determined, and it has been agreed to allow the symbols to stand for these relative weights, an equation or formula making use of the symbols becomes a statement of a definite numerical fact, and calculations can be based on it.

Chemical equations not algebraic. Although chemical equations are quantitative, it must be clearly understood that they are not algebraic. A glance at the equations

will show at once that they are true. The equations

are equally true in an algebraic sense, but experiment shows that only the first is true chemically, for iron oxide (FeO)[Pg 75] cannot be directly decomposed into iron and oxygen. Only such equations as have been found by careful experiment to express a real chemical transformation, true both for the kinds of substances as well as for the weights, have any value.

Chemical formulas and equations, therefore, are a concise way of representing qualitatively and quantitatively facts which have been found by experiment to be true in reference to the composition of substances and the changes which they undergo.

Formulas representing water of crystallization. An examination of substances containing water of crystallization has shown that in every case the water is present in such proportion by weight as can readily be represented by a formula. For example, copper sulphate (CuSO₄) and water combine in the ratio of 1 molecule of the sulphate to 5 of water; calcium sulphate (CaSO₄) and water combine in the ratio 1: 2 to form gypsum. These facts are expressed by writing the formulas for the two substances with a period between them. Thus the formula for crystallized copper sulphate is CuSO₄·5H₂O; that of gypsum is CaSO₄·2H₂O.

Heat of reaction. Attention has frequently been directed to the fact that chemical changes are usually accompanied by heat changes. In general it has been found that in every chemical action heat is either absorbed or given off. By adopting a suitable unit for the measurement of heat, the heat change during a chemical reaction can be expressed in the equation for the reaction.

Heat cannot be measured by the use of a thermometer alone, since the thermometer measures the intensity of heat, not its quantity. The easiest way to measure a quantity of heat is to note how warm it will make a definite amount of[Pg 76] a given substance chosen as a standard. Water has 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 raise the temperature of one gram of water one degree.

By means of this unit it is easy to indicate the heat changes in a given chemical reaction. The equation

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.

means that an expenditure of 19,000 cal. is required to cause 12 g. of carbon to unite with 64.12 g. of sulphur to form 76.12 g. of carbon disulphide. In these equations it will be noted that the symbols stand for as many grams of the substance as there are units in the weights of the atoms represented by the symbols. This is always understood to be the 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 take place. The equation

does not tell us that it is necessary to keep the mercuric oxide at a high temperature in order that the decomposition may go on. The equation

in no way indicates the fact that the hydrochloric acid must be dissolved in water before it will act upon the zinc. From the equation

[Pg 77]

it would not be suspected that the two gases hydrogen and chlorine will unite instantly in the sunlight, but will stand mixed in the dark a long time without change. It will therefore be necessary to pay much attention to the details of the conditions under which a given reaction occurs, as well as to the expression of the reaction in the form of an equation.

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 compositions as given in Chapters II and III.

2. Determine the percentage of copper, sulphur, oxygen, and water in copper sulphate crystals. What weight of water can be obtained from 150 g. of this substance?

3. What weight of zinc can be dissolved in 10 g. of sulphuric acid? How much zinc sulphate will be formed?

4. How many liters of hydrogen measured under standard conditions can be obtained from the action of 8 g. of iron on 10 g. of sulphuric acid? How much iron sulphate (FeSO₄) will be formed?

5. 10 g. of zinc were used in the preparation of hydrogen; what weight of iron will be required to prepare an equal volume?

6. How many grams of barium dioxide will be required to prepare 1 kg. of common hydrogen dioxide solution? What weight of barium sulphate will be formed at the same time?

7. What weight of the compound Mn₃O₄ will be formed by strongly heating 25 g. of manganese dioxide? What volume of oxygen will be given off at the same time, measured under standard conditions?

8. (a) What is the weight of 100 l. of hydrogen measured in a laboratory in which the temperature is 20° and pressure 750 mm.? (b) What weight of sulphuric acid is necessary to prepare this amount of hydrogen? (c) The density of sulphuric acid is 1.84. Express the acid required in (b) in cubic centimeters.

9. What weight of potassium chlorate is necessary to furnish sufficient oxygen to fill four 200 cc. bottles in your laboratory (the gas to be collected over water)?

[Pg 78]

CHAPTER VII

NITROGEN AND THE RARE ELEMENTS: ARGON, HELIUM, NEON, KRYPTON, XENON

Historical. Nitrogen was discovered by the English chemist Rutherford in 1772. A little later Scheele showed it to be a constituent of air, and Lavoisier gave it the name azote, signifying that it would not support life. The name nitrogen was afterwards given it because of its presence in saltpeter or niter. The term azote and symbol Az are still retained by the French chemists.

Occurrence. Air is composed principally of oxygen and nitrogen in the free state, about 78 parts by volume out of every 100 parts being nitrogen. Nitrogen also occurs in nature in the form of potassium nitrate (KNO₃)—commonly called saltpeter or niter—as well as in sodium nitrate (NaNO₃). Nitrogen is also an essential constituent of all living organisms; for example, the human body contains about 2.4% of nitrogen.

Preparation from air. Nitrogen can be prepared from air by the action of some substance which will combine with the oxygen, leaving the nitrogen free. Such a substance must be chosen, however, as will combine with the oxygen to form a product which is not a gas, and which can be readily separated from the nitrogen. The substances most commonly used for this purpose are phosphorus and copper.

1. By the action of phosphorus. The method used for the preparation of nitrogen by the action of phosphorus is as follows:[Pg 79]

The phosphorus is placed in a little porcelain dish, supported on a cork and floated on water (Fig. 26). It is then ignited by contact with a hot wire, and immediately a bell jar or bottle is brought over it so as to confine a portion of the air. The phosphorus combines with the oxygen to form an oxide of phosphorus, known as phosphorus pentoxide. This is a white solid which floats about in the bell jar, but in a short time it is all absorbed by the water, leaving the nitrogen. The withdrawal of the oxygen is indicated by the rising of the water in the bell jar.

Fig. 26

2. By the action of copper. The oxygen present in the air may also be removed by passing air slowly through a heated tube containing copper. The copper combines with the oxygen to form copper oxide, which is a solid. The nitrogen passes on and may be collected over water.

Preparation from compounds of nitrogen. Pure nitrogen may be obtained from certain compounds of the element. Thus, if heat is applied to the compound ammonium nitrite (NH₄NO₂), the change represented in the following equation takes place:

Physical properties. Nitrogen is similar to oxygen and hydrogen in that it is a colorless, odorless, and tasteless gas. One liter of nitrogen weighs 1.2501 g. It is almost insoluble in water. It can be obtained in the form of a colorless liquid having a boiling point of -195° at ordinary pressure. At -214° it solidifies.

Chemical properties. Nitrogen is characterized by its inertness. It is neither combustible nor a supporter of combustion. At ordinary temperatures it will not combine directly with any of the elements except under rare conditions. At higher temperatures it combines with magnesium, lithium, titanium, and a number of other elements. The compounds formed are called nitrides, just as compounds of an element with oxygen are called oxides. When it is mixed with oxygen and subjected to the action of electric sparks, the two gases slowly combine forming oxides of nitrogen. A mixture of nitrogen and hydrogen when treated similarly forms ammonia, a gaseous compound of nitrogen and hydrogen. Since we are constantly inhaling nitrogen, it is evident that it is not poisonous. Nevertheless life would be impossible in an atmosphere of pure nitrogen on account of the exclusion of the necessary oxygen.

[Pg 82]

EXERCISES

1. How could you distinguish between oxygen, hydrogen, and nitrogen?

2. Calculate the relative weights of nitrogen and oxygen; of nitrogen and hydrogen.

3. In the preparation of nitrogen from the air, how would hydrogen do as a substance for the removal of the oxygen?

4. What weight of nitrogen can be obtained from 10 l. of air measured under the conditions of temperature and pressure which prevail in your laboratory?

5. How many grams of ammonium nitrite are necessary in the preparation of 20 l. of nitrogen measured over water under the conditions of temperature and pressure which prevail in your laboratory?

6. If 10 l. of air, measured under standard conditions, is passed over 100 g. of hot copper, how much will the copper gain in weight?

WILLIAM RAMSAY (Scotch) (1855-)

Has made many studies in the physical properties of substances; discovered helium; together with Lord Rayleigh and others he discovered argon, krypton, xenon, and neon; has contributed largely to the knowledge of radio-active substances, showing that radium gradually gives rise to helium; professor at University College, London

[Pg 83]

CHAPTER VIII

THE ATMOSPHERE

Atmosphere and air. The term atmosphere is applied to the gaseous envelope surrounding the earth. The term air is generally applied to a limited portion of this envelope, although the two words are often used interchangeably. Many references have already been made to the composition and properties of the atmosphere. These statements must now be collected and discussed somewhat more in detail.

Air formerly regarded as an element. Like water, air was at first regarded as elementary in character. Near the close of the eighteenth century Scheele, Priestley, and Lavoisier showed by their experiments that it is a mixture of at least two gases,—those which we now call oxygen and nitrogen. By burning substances in an inclosed volume of air and noting the contraction in volume due to the removal of the oxygen, they were able to determine with some accuracy the relative volumes of oxygen and nitrogen present in the air.

The constituents of the atmosphere. The constituents of the atmosphere may be divided into two general groups: those which are essential to life and those which are not essential.

1. Constituents essential to life. In addition to oxygen and nitrogen at least two other substances, namely, carbon dioxide and water vapor, must be present in the atmosphere in order that life may exist. The former of these is a[Pg 84] gaseous compound of carbon and oxygen having the formula CO₂. Its properties will be discussed in detail in the chapter on the compounds of carbon. Its presence in the air may be shown by causing the air to bubble through a solution of calcium hydroxide (Ca(OH)₂), commonly called lime water. The carbon dioxide combines with the calcium hydroxide in accordance with the following equation:

The resulting calcium carbonate (CaCO₃) is insoluble in water and separates in the form of a white powder, which causes the solution to appear milky.

The presence of water vapor is readily shown by its condensation on cold objects as well as by the fact that a bit of calcium chloride when exposed to the air becomes moist, and may even dissolve in the water absorbed from the air.

2. Constituents not essential to life. In addition to the essential constituents, the air contains small percentages of various other gases, the presence of which so far as is known is not essential to life. This list includes the rare elements, argon, helium, neon, krypton, and xenon; also hydrogen, ammonia, hydrogen dioxide, and probably ozone. Certain minute forms of life (germs) are also present, the decay of organic 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[Pg 85] order, the Leguminosæ, 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.

Fig. 27

The quantitative analysis of air. A number of different methods have been devised for the determination of the percentages of the constituents present in the atmosphere. Among these are the following.

1. Determination of oxygen. (1) The oxygen is withdrawn from a measured volume of air inclosed in a tube, by means of phosphorus.

Fig. 28

[Pg 86]

(2) The oxygen may also be estimated by passing a measured volume of air through a tube containing copper heated to a high temperature. The oxygen 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 the oxygen in the volume of air taken.

(3) A more accurate method is the following. A eudiometer tube is filled with mercury and inverted in a vessel of the same liquid. A convenient amount of air is then introduced into the tube and its volume accurately noted. There is then introduced more than sufficient hydrogen to combine with the oxygen present in the inclosed air, and the volume is again accurately noted. The mixture is then exploded by an electric spark, and the volume is once more taken. By subtracting this volume from the total volume of the air and hydrogen there is obtained the contraction in volume due to the union of the oxygen and hydrogen. The volume occupied by the water formed by the union of the two gases is so small that it may be disregarded in the calculation. Since oxygen and hydrogen combine in the ratio 1: 2 by volume, it is evident that the contraction in volume due to the combination is equal to the volume occupied by the oxygen in the air contained in the tube, plus twice this volume of hydrogen. In other words, one third of the total contraction is equal to the volume occupied by the oxygen in the inclosed air. The following example will make this clear:

[Pg 87]

All these methods agree in showing that 100 volumes of dry air contain approximately 21 volumes of oxygen.

2. Determination of nitrogen. If the gas left after the removal of oxygen from a portion of air is passed over heated magnesium, the nitrogen is withdrawn, argon and the other rare elements being left. It may thus be shown that of the 79 volumes of gas left after the removal of the oxygen from 100 volumes of air, approximately 78 are nitrogen and 0.93 argon. The other elements are present in such small quantities that they may be neglected.

3. Determination of carbon dioxide. The percentage of carbon dioxide in any given volume of air may be determined by passing the air over calcium hydroxide or some other compound which will combine with the carbon dioxide. The increase in the weight of the hydroxide equals the weight of the carbon dioxide absorbed. The amount present in the open normal air is from 3 to 4 parts by volume in 10,000 volumes of air, or about 0.04%.

4. Determination of water vapor. The water vapor present in a given volume of air may be determined by passing the air over calcium chloride (or some other compound which has a strong affinity for water), and noting the increase in the weight of the chloride. The amount present varies not only with the locality, but there is a wide variation from day to day in the same locality because of the winds and changes in temperature.

Processes affecting the composition of the air. The most important of these processes are the following.

1. Respiration. In the process of respiration some of the oxygen in the inhaled air is absorbed by the blood and carried to all parts of the body, where it combines with the carbon of the worn-out tissues. The products of oxidation[Pg 88] are carried back to the lungs and exhaled in the form of carbon dioxide. The amount exhaled by an adult averages about 20 l. per hour. Hence in a poorly ventilated room occupied by a number of people the amount of carbon dioxide rapidly increases. While this gas is not poisonous unless present in large amounts, nevertheless air containing more than 15 parts in 10,000 is not fit for respiration.

2. Combustion. All of the ordinary forms of fuel contain large percentages of carbon. On burning, this carbon combines with oxygen in the air, forming carbon dioxide. Combustion and respiration, therefore, tend to diminish the amount of oxygen in the air and to increase the amount of carbon dioxide.

3. Action of plants. Plants have the power, when in the sunlight, of absorbing carbon dioxide from the air, retaining the carbon and returning at least a portion of the oxygen to the air. It will be observed that these changes are just the opposite of those brought about by the processes of respiration and combustion.

The composition of the air is constant. Notwithstanding the changes constantly taking place which tend to alter the composition of the air, the results of a great many analyses of air collected in the open fields show that the percentages of oxygen and nitrogen as well as of carbon dioxide are very nearly constant. Indeed, so constant are the percentages of oxygen and nitrogen that the question has arisen, whether these two elements are not combined in the air, forming a definite chemical compound. That the two are not combined but are simply mixed together can be shown in a number of ways, among which are the following.[Pg 90]

1. When air dissolves in water it has been found that the ratio of oxygen to nitrogen in the dissolved air is no longer 21: 78, but more nearly 35: 65. If it were a chemical compound, the ratio of oxygen to nitrogen would not be changed by solution in water.

2. A chemical compound in the form of a liquid has a definite boiling point. Water, for example, boils at 100°. Moreover the steam which is thus formed has the same composition as the water. The boiling point of liquid air, on the other hand, gradually rises as the liquid boils, the nitrogen escaping first followed by the oxygen. If the two were combined, they would pass off together in the ratio in which they are found in the air.

Why the air has a constant composition. If air is a mixture and changes are constantly taking place which tend to modify its composition, how, then, do we account for the constancy of composition which the analyses reveal? This is explained by several facts. (1) The changes which are caused 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 so small compared with the total amount of air that they cannot be readily detected. (4) Finally it must be noted that only air collected in the open fields shows this constancy in composition. The air in a poorly ventilated room occupied by a number of people rapidly changes in composition.

The properties of the air. Inasmuch as air is composed principally of a mixture of oxygen and nitrogen, which elements have already been discussed, its properties may be inferred largely from those of the two gases.[Pg 91] One liter weighs 1.2923 g. It is thus 14.38 times as heavy as hydrogen. At the sea level it exerts an average pressure sufficient to sustain a column of mercury 760 mm. in height. This is taken as the standard pressure in determining the volumes of gases as well as the boiling points of liquids. Water may be made to boil at any temperature between 0° and considerably above 100° by simply varying the pressure. It is only when the pressure upon it is equal to the normal pressure of the atmosphere at the sea level, as indicated by a barometric reading of 760 mm., that it boils at 100°.

Preparation of liquid air. Attention has been called to the fact that both oxygen and nitrogen can be obtained in the liquid state by strongly cooling the gases and applying great pressure to them. Since air is largely a mixture of these two gases, it can be liquefied by the same methods.

Fig. 29

The Dewar bulb. It is not possible to preserve air in the liquid state in a closed vessel, on account of the enormous pressure exerted by it in its tendency to pass into the gaseous state. It may however be preserved for some hours or even days before it will completely evaporate, by simply placing it in an open vessel surrounded by a nonconducting material. The most efficient vessel for this purpose is the Dewar bulb shown in Fig. 29.[Pg 92] The air is withdrawn from the space between the two walls, thus making it nonconducting.

Properties and uses of liquid air. When first prepared, liquid air is cloudy because of the presence of particles of solid carbon dioxide. These may be filtered off, leaving a liquid of slightly bluish color. It begins to boil at about -190°, the nitrogen passing off first, gradually followed by the oxygen, the last portions being nearly pure oxygen. To a certain extent oxygen is now prepared in this way for commercial purposes.

The extremely low temperature of liquid air may be inferred from the fact that mercury when cooled by it is frozen to a mass so hard that it may be used for driving nails.

Liquid air is used in the preparation of oxygen and as a cooling agent in the study of the properties of matter at low temperatures. It has thus been found that elements at extremely low temperatures largely lose their chemical activity.

EXERCISES

1. When oxygen and nitrogen are mixed in the proportion in which they exist in the atmosphere, heat is neither evolved nor absorbed by the process. What important point does this suggest?

2. What essential constituent of the air is found in larger amount in manufacturing districts than in the open country?

3. Can you suggest any reason why the growth of clover in a field improves the soil?

4. Why are the inner walls of a Dewar bulb sometimes coated with a film of silver?

5. To what is the blue color of liquid air due? Does this color increase in intensity on standing?

6. When ice is placed in a vessel containing liquid air, the latter boils violently. Explain.[Pg 93]

7. Taking the volumes of the oxygen and nitrogen in 100 volumes of air as 21 and 78 respectively, calculate the percentages of these elements present by weight.

8. Would combustion be more intense in liquid air than in the gaseous substance?

9. A tube containing calcium chloride was found to weigh 30.1293 g. A volume of air which weighed 15.2134 g. was passed through, after which the weight of the tube was found to be 30.3405 g. What was the percentage amount of moisture present in the air?

10. 10 l. of air measured at 20° and 740 mm. passed through lime water caused the precipitation of 0.0102 g. of CaCO₃. Find the number of volumes of carbon dioxide in 10,000 volumes of the air.

[Pg 94]

CHAPTER IX

SOLUTIONS

Definitions. When a substance disappears in a liquid in such a way as to thoroughly mix with it and to be lost to sight as an individual body, the resulting liquid is called a solution. The liquid in which the substance dissolves is called the solvent, while the dissolved substance is called the solute.

Classes of solutions. Matter in any one of its physical states may dissolve in a liquid, so that we may have solutions of gases, of liquids, and of solids. Solutions of liquids in liquids are not often mentioned in the following pages, but the other two classes will become very familiar in the course of our study, and deserve special attention.

SOLUTION OF GASES IN LIQUIDS

Fig. 30

It has already been stated that oxygen, hydrogen, and nitrogen are slightly soluble in water. Accurate study has led to the conclusion that all gases are soluble to some extent not only in water but in many other liquids. The amount of a gas which will dissolve in a liquid depends upon a number of conditions, and these can best be understood by[Pg 95] supposing a vessel B (Fig. 30), to be filled with the gas and inverted over the liquid. Under these circumstances the gas cannot escape or become mixed with another gas.

Circumstances affecting the solubility of gases. A number of circumstances affect the solubility of a gas in a liquid.

1. Nature of the gas. Other conditions being equal, each gas has its own peculiar solubility, just as it has its own special taste or odor. The solubility of gases varies between wide limits, as will be seen from the following table, but as a rule a given volume of a liquid will not dissolve 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:

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 gas takes place.

2. Nature of the liquid. The character of the liquid has much influence upon the solubility of a gas. Water, alcohol, and ether have each its own peculiar solvent power. From the solubility of a gas in water, no prediction can be made as to its solubility in other liquids.

3. Influence of pressure. It has been found that the weight of gas which dissolves in a given case is proportional to the pressure exerted upon the gas. If the[Pg 96] pressure is doubled, the weight of gas going into solution is doubled; if the pressure is diminished to one half of its original value, half of the dissolved gas will escape. Under high pressure, large quantities of gas can be dissolved in a liquid, and when the pressure is removed the gas escapes, causing the liquid to foam or effervesce.

4. Influence of temperature. In general, the lower the temperature of the liquid, the larger the quantity of gas which it can dissolve. 1000 volumes 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 expelled from a liquid by boiling the solution, some cannot. For example, it is not possible to expel hydrochloric acid gas completely from its solution by boiling.

SOLUTION OF SOLIDS IN LIQUIDS

This is the most familiar class of solutions, since in the laboratory substances are much more frequently used in the form of solutions than in the solid state.

Circumstances affecting the solubility of a solid. The solubility of a solid in a liquid depends upon several factors.

1. Nature of the solid. Other conditions being the same, solids vary greatly in their solubility in liquids. This is illustrated in the following table:

Table of Solubility of Solids at 18°

[Pg 97]

No solids are absolutely insoluble, but the amount dissolved may be so small as to be of no significance for most purposes. Thus barium sulphate, one of the most insoluble of common substances, dissolves in water to the extent of 1 part in 400,000.

2. Nature of the solvent. Liquids vary much in their power to dissolve solids. Some are said to be good solvents, since they dissolve a great variety of substances and considerable quantities of them. Others have small solvent power, dissolving few substances, and those to a slight extent only. Broadly speaking, water is the most general solvent, and alcohol is perhaps second in solvent power.

3. Temperature. The weight of a solid which a given liquid can dissolve varies with the temperature. Usually it increases rapidly as the temperature rises, so that the boiling liquid dissolves several times the weight which the cold liquid will dissolve. In some instances, as in the case of common salt dissolved in water, the temperature has little influence upon the solubility, and a few solids are more soluble in cold water than in hot. The following examples will serve as illustrations:

Table of Solubility at 0° and at 100°

100 cc. of water will dissolve:

Saturated solutions. A liquid will not dissolve an unlimited quantity of a solid. On adding the solid to the liquid in small portions at[Pg 98] a time, it will be found that a point is reached at which the liquid will not dissolve more of the solid at that temperature. The solid and the solution remain in contact with each other unchanged. This condition may be described by saying that they are in equilibrium with each other. A solution is said to be saturated when it remains unchanged in concentration in contact with some of the solid. The weight of the solid which will completely saturate a definite volume of a liquid at a given temperature is called the solubility of the substance at that temperature.

Supersaturated solutions. When a solution, saturated at a given temperature, is allowed to cool it sometimes happens that no solid crystallizes out. This is very likely to occur when the vessel used is perfectly smooth and the solution is not disturbed in any way. Such a solution is said to be supersaturated. That this condition is unstable can be shown by adding a crystal of the solid to the solution. All of the solid in excess of the quantity required to saturate the solution at this temperature will at once crystallize out, leaving the solution saturated. Supersaturation may also be overcome in many cases by vigorously shaking or stirring the solution.

General physical properties of solutions. A few general statements may be 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 every part of the solution has the same concentration. The process goes on very slowly unless hastened by stirring or shaking the solution. Thus, if a few crystals of a highly colored substance such as copper sulphate are placed in the bottom of a tall vessel full of water, it will take weeks for the solution to become uniformly colored.

2. Boiling points of solutions. The boiling point of a liquid is raised by the presence of a substance dissolved in it. In general the extent to which the boiling point of a solvent is raised by a given substance is proportional to the[Pg 99] concentration of the solution, that is, to the weight of the substance dissolved in a definite weight of the solvent.

3. Freezing points of solutions. A solution freezes at a lower temperature than the pure solvent. The lowering of the freezing point obeys the same law which holds for the raising of the boiling point: the extent 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 the electric current. If, however, certain substances such as common salt are dissolved in the water, the resulting solutions are found to be conductors of electricity. Such solutions are called electrolytes. When the current passes through an electrolyte some chemical change always takes place. This change is called electrolysis.

Fig. 31

The general method used in the electrolysis of a solution is illustrated in Fig. 31. The vessel D contains the electrolyte. Two plates or rods, A and B, made of suitable material, are connected with the wires from a battery (or dynamo) and dipped into the electrolyte, as shown in the figure. These plates or rods are called electrodes. The electrode connected with the zinc plate of the battery is the negative electrode or cathode, while that connected with the carbon plate is the positive electrode or anode.

Theory of electrolytic dissociation. The facts which have just been described in connection with solutions, together with many others, have led chemists to adopt a theory of solutions called the theory of electrolytic dissociation. The main assumptions in this theory are the following.[Pg 100]

1. Formation of ions. Many compounds when dissolved in water undergo an important change. A portion of their molecules fall apart, or dissociate, into two or more parts, called ions. Thus sodium nitrate (NaNO₃) dissociates into the ions Na and NO₃; sodium chloride, into the ions Na and Cl. These ions are free to move about in the solution independently of each other like independent molecules, and for this reason were given the name ion, which signifies a wanderer.

2. The electrical charge of ions. Each ion carries a heavy electrical charge, and in this respect differs from an atom or molecule. It is evident that the sodium in the form of an ion must differ in some important way from ordinary sodium, for sodium ions, formed from sodium nitrate, give no visible evidence of their presence in water, whereas metallic sodium at once decomposes the water. The electrical charge, therefore, greatly modifies the usual chemical properties of the element.

3. The positive charges equal the negative charges. The ions formed by the dissociation of any molecule are of two kinds. One kind is charged with positive electricity and the other with negative electricity; moreover the sum of all the positive charges is always equal to the sum of all the negative charges. The solution as a whole is therefore electrically neutral. If we represent dissociation by the usual chemical equations, with the electrical charges indicated by + and - signs following the symbols, the dissociation of sodium chloride molecules is represented thus:

The positive charge on each sodium ion exactly equals the negative charge on each chlorine ion.[Pg 101] Sodium sulphate dissociates, as shown in the equation

Here the positive charge on the two sodium ions equals the double negative charge on the SO₄ ion.

4. Not all compounds dissociate. Only those compounds dissociate whose solutions form electrolytes. Thus salt dissociates when dissolved in water, the resulting solution being an electrolyte. Sugar, on the other hand, does not dissociate and its solution is not a conductor of the electric current.

5. Extent of dissociation differs in different liquids. While compounds most readily undergo dissociation in water, yet dissociation often occurs to a limited extent when solution takes place in liquids other than water. In the discussion of solutions it will be understood that 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 explanation of the properties of solutions. Let us now see if the theory is in harmony with certain of these properties.

The theory of electrolytic dissociation and the boiling and freezing points of solutions. We have seen that the boiling point of a solution of a substance is raised in proportion to the concentration of the dissolved substance. This is but another way of saying that the change in the boiling point of the solution is proportional to the number of molecules of the dissolved substance present in the solution.

It has been found, however, that in the case of electrolytes the boiling point is raised more than it should be to[Pg 102] conform to this law. If the solute dissociates into ions, the reason for this becomes clear. Each ion has the same effect on the boiling point as a molecule, and since their number is greater than the number of molecules from which they were formed, the effect on the boiling point is abnormally great.

In a similar way, the theory furnishes an explanation of the abnormal lowering of the freezing point of electrolytes.

The theory of electrolytic dissociation and electrolysis. The changes taking place during electrolysis harmonize very completely with the theory of dissociation. This will become clear from a study of the following examples.

Fig. 32

1. Electrolysis of sodium chloride. Fig. 32 represents a vessel in which the electrolyte is a solution of sodium chloride (NaCl). According to the dissociation theory the molecules of sodium chloride dissociate into the ions Na⁺ and Cl^-. The Na⁺ ions are attracted to the cathode owing to its large negative charge. On coming into contact with the cathode, the Na⁺ ions give up their positive charge and are then ordinary sodium atoms. They immediately decompose the water according to the equation

and hydrogen is evolved about the cathode.

The chlorine ions on being discharged at the anode in similar manner may either be given off as chlorine gas, or may attack the water, as represented in the equation

2. Electrolysis of water. The reason for the addition of sulphuric acid to water in the preparation of oxygen and hydrogen by electrolysis will now be clear. Water itself is not an electrolyte to an appreciable extent; that is, it does not form enough ions to carry a current. Sulphuric acid dissolved in water is an electrolyte, and dissociates into the ions 2 H⁺ and SO₄^—. In the process of electrolysis of the solution, the hydrogen ions travel to the cathode, and on being discharged escape as hydrogen gas. The SO₄ ions, when discharged at the anode, act upon water, setting free oxygen and once more forming sulphuric acid:

The sulphuric acid can again dissociate and the process repeat itself as long as any water is left. Hence the hydrogen and oxygen set free in the electrolysis of water really come directly from the acid but indirectly from the water.

3. Electrolysis of sodium sulphate. In a similar way, sodium sulphate (Na₂SO₄), when in solution, gives the ions 2 Na⁺ and SO₄^—. On being discharged, the sodium atoms decompose water about the cathode, as in the case of sodium chloride, while the SO₄ ions when discharged at the anode decompose the water, as represented in the equation

Fig. 33

That new substances are formed at the cathode and anode may be shown in the following way. A U-tube, such as is represented in Fig. 33, is partially filled with a solution of sodium sulphate, and the liquid in one arm is colored with red litmus, that in the other[Pg 104] 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 the current to pass, the blue solution turns red, while the red solution turns blue. These are exactly the changes which would take place if sodium hydroxide and sulphuric acid were to be set free at the electrodes, as required by the theory.

The properties of electrolytes depend upon the ions present. When a substance capable of dissociating into ions is dissolved in water, the properties of the solution will depend upon two factors: (1) the ions formed from the substance; (2) the undissociated molecules. Since the ions are usually more active chemically than the molecules, most of the chemical properties of an electrolyte are due to the ions rather than to the molecules.

The solutions of any two substances which give the same ion will have certain properties in common. Thus all solutions containing the copper ion (Cu⁺⁺) are blue, unless the color is modified by the presence of ions 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 ice and salt produce a lower temperature than ice alone?

4. Why does shaking or stirring make a solid dissolve more rapidly in a liquid?

5. By experiment it was found that a certain volume of water was saturated at 100° with 114 g. of potassium nitrate. On cooling to 0° a portion of the substance crystallized. (a) How many grams of the substance remained in solution? (b) What was the strength[Pg 105] of the solution 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 solution evaporated to dryness; what weight of solid was left? (b) 10 g. of zinc were dissolved in hydrochloric acid and the solution evaporated to dryness; 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 (CuSO₄); this solution is blue. (b) In the same manner, write one for sodium sulphate; this solution is colorless. (c) How would you account for the color of the copper sulphate solution?

9. (a) As in the preceding exercise, write a simple equation for a dilute solution of copper chloride (CuCl₂); this solution is blue. (b) In the same manner, write one for sodium chloride; this solution is colorless. To what is the blue color due?

10. What component is present in concentrated sulphuric acid that is almost wanting in very dilute sulphuric acid?

11. Why will vegetables cook faster when boiled in strong salt water than when boiled in pure water?

12. How do you explain the foaming of soda water?

[Pg 106]

CHAPTER X

ACIDS, BASES, AND SALTS; NEUTRALIZATION

Acids, bases, and salts. The three classes of compounds known respectively as acids, bases, and salts include the great majority of the 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 in detail in the appropriate places, but a few representatives of each class will be described in this chapter with special reference to the common properties in accordance with which they are classified.

The familiar acids. Hydrochloric acidis a gas composed of hydrogen and chlorine, and has the formula HCl. The substance is very soluble in water, and it is this solution which is usually called hydrochloric acid. Nitric acid is a liquid composed of hydrogen, nitrogen, and oxygen, having the formula HNO₃. As sold commercially it is mixed with about 32% of water. Sulphuric acid, whose composition is represented by the formula H₂SO₄, is an oily liquid nearly twice as heavy as water, and is commonly called oil of vitriol.

Characteristics of acids. (1) All acids contain hydrogen. (2) When dissolved in water the molecules of the acid dissociate into two kinds of ions. One of these is always hydrogen and is the cation (+), while the other consists of the remainder of the molecule and is the anion (-). (3) The solution tastes sour. (4) It has the power to change the[Pg 107] color of certain substances called indicators. Thus blue litmus is changed to red, and yellow methyl orange is changed to red. Since all acids produce hydrogen cations, while the anions of each are different, the properties which all acids have in common when in solution, such as taste and action on indicators, must be attributed to the hydrogen ions.

DEFINITION: An acid is a substance which produces hydrogen ions when dissolved in water or other dissociating liquids.

Undissociated acids. When acids are perfectly free from water, or are dissolved in liquids like benzene which do not have the power of dissociating them into ions, they should have no real acid properties. This is found to be the case. Under these circumstances they do not affect the color of indicators or have any of the properties characteristic of acids.

The familiar bases. The bases most used in the laboratory are sodium hydroxide (NaOH), potassium hydroxide (KOH), and calcium hydroxide (Ca(OH)₂). These are white solids, soluble in water, the latter sparingly so. Some bases are very difficultly soluble in water. The very soluble ones with most pronounced basic properties are sometimes called the alkalis.

Characteristics of bases. (1) All bases contain hydrogen and oxygen. (2) When dissolved in water the molecules of the base dissociate into two kinds of ions. One of these is always composed of oxygen and hydrogen and is the anion. It has the formula OH and is called the hydroxyl ion. The remainder of the molecule, which usually consists of a single atom, is the cation. (3) The solution of a base has[Pg 108] a soapy feel and a brackish taste. (4) It reverses the color change produced in indicators by acids, turning red litmus blue, and red methyl orange yellow. Since all bases produce hydroxyl anions, while the cations of each are different, the properties which all bases have in common when in solution must be due to the hydroxyl ions.

DEFINITION: A base is a substance which produces hydroxyl ions when dissolved in water or other dissociating liquids.

Undissociated bases. Bases, in the absence of water or when dissolved in liquids which do not dissociate them, should have none of the properties characteristic of this class of substances. This has been found to be the case. For example, they have no effect upon indicators under these circumstances.

Neutralization. When an acid and a base are brought together in solution in proper proportion, the characteristic properties of each disappear. The solution tastes neither sour nor brackish; it has no effect upon indicators. There can therefore be neither hydrogen nor hydroxyl ions present in the solution. A study of reactions of this kind has shown that the hydrogen ions of the acid combine with the hydroxyl ions of the base to form molecules of water, water being a substance which is not appreciably dissociated into ions. This action of an acid on a base is called neutralization. The following equations express the neutralization of the three acids by three bases, water being formed in each case.

[Pg 109]

DEFINITION: Neutralization consists in the union of the hydrogen ion of an acid with the hydroxyl ion of a base to form water.

Salts. It will be noticed that in neutralization the anion of the acid and the cation of the base are not changed. If, however, the water is expelled by evaporation, these two ions slowly unite, and when the water becomes saturated with the substance so produced, it separates in the form of a solid called a salt.

DEFINITION: A salt is a substance formed by the union of the anion of an acid with the cation of a base.

Characteristics of salts. (1) From the definition of a salt it will be seen that there is no element or group of elements which characterize salts. (2) Salts as a class have no peculiar taste. (3) In the absence of all other substances they are without action on indicators. (4) When dissolved in water they form two kinds of ions.

Neutralization a quantitative act. Since neutralization is a definite chemical act, each acid will require a perfectly definite weight of each base for its neutralization. For[Pg 110] example, a given weight of sulphuric acid will always require a definite weight of sodium hydroxide, in accordance with the equation

Fig. 34

Extent of dissociation. The question will naturally arise, When an acid, base, or salt dissolves in water, do all the molecules dissociate into ions, or only a part of them? The experiments by which this question can be answered cannot be described here. It has been found, however, that only a fraction of the molecules dissociate. The percentage which will dissociate in a given case depends upon several conditions, the chief of which are: (1) The concentration of the solution. In concentrated solutions only a very small[Pg 111] percentage of dissociation occurs. As the solution is diluted the percentage increases, and in very dilute solutions it may be very large, though it is never complete in any ordinary solution. (2) The nature of the dissolved compound. At equal concentrations substances differ much among themselves in the percentage of dissociation. The great majority of salts are about equally dissociated. Acids and bases, on the contrary, show great differences. Some are freely dissociated, while others are dissociated to but a slight extent.

Strength of acids and bases. Since acid and basic properties are due to hydrogen and hydroxyl ions respectively, the acid or base which will produce the greatest percentage of these ions at a given concentration must be regarded as the strongest representative of its class. The acids and bases described in the foregoing paragraphs are all quite strong. In 10% solutions they are dissociated to about 50%, and this is also approximately the extent to which most salts are dissociated at this same concentration.

Methods of expressing reactions between compounds in solution. Chemical equations representing reactions between substances in solution may represent the details of the reaction, or they may simply indicate the final products formed. In the latter case the formation of ions is not indicated. Thus, if we wish to call attention to the details of the reaction between sodium hydroxide and hydrochloric acid in solution, the equation is written as follows:

On the other hand, if we wish simply to represent the final products formed, the following is used.

Both of these methods will therefore be used:

Radicals. It has been emphasized that the hydroxyl group (OH) always [Pg 113]forms the anion of a base, while the group NO₃ forms the anion of nitric acid and sodium nitrate; the group SO₄, the anion of sulphuric acid and calcium sulphate. A group of elements which in this way constitutes a part of a molecule, acting as a unit in a chemical change, or forming ions in solution, is called a radical. Some of these radicals have been given special names, the names signifying the elements present in the radical. Thus we have the hydroxyl radical (OH) and the nitrate radical (NO₃).

DEFINITION: A radical is a group of elements forming part of a molecule, and acting as a unit in chemical reactions.

Names of acids, bases, and salts. Since acids, bases, and salts are so intimately related to each other, it is very advantageous to give names to the three classes in accordance with some fixed system. The system universally adopted is as follows:

Naming of bases. All bases are called hydroxides. They are distinguished from each other by prefixing the name of the element which is 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 the acid consists of two elements or three.

1. Binary acids. Acids containing only one element in addition to hydrogen are called binary acids. They are given names consisting of the prefix hydro-, the name of the second element present, and the termination -ic. Examples: hydrochloric acid (HCl); hydrosulphuric acid (H₂S).

2. Ternary acids. In addition to the two elements present in binary acids, the great majority of acids also contain oxygen. They therefore consist of three elements and[Pg 114] are called ternary acids. It usually happens that the same three elements can unite in different proportions to make several different acids. The most familiar one of these is given a name ending in the suffix -ic, while the one with less oxygen is given a similar name, but ending in the suffix -ous. Examples: nitric acid (HNO₃); nitrous acid (HNO₂). In cases where more than two acids are known, use is made of prefixes in addition to the two suffixes -icand -ous. Thus the prefix per- signifies an acid still richer in oxygen; the prefix hypo- signifies one with less oxygen.

Naming of salts. A salt derived from a binary acid is given a name consisting of the names of the two elements composing it, with the termination -ide. Example: sodium chloride (NaCl). All other binary compounds are named in the same way.

A salt of a ternary acid is named in accordance with the acid from which it is derived. A ternary acid with the termination -ic gives a salt with the name ending in -ate, while an acid with termination -ous gives a salt with the name ending in -ite. The following table will make the application of these principles clear:

[Pg 115]

EXERCISES

1. 25 cc. of a solution containing 40 g. of sodium hydroxide per liter was found to neutralize 25 cc. of a solution of hydrochloric acid. What was 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 of each substance had been used?

3. A solution contains 18 g. of hydrochloric acid per 100 cc. It required 25 cc. of this solution to neutralize 30 cc. of a solution of sodium hydroxide. What was the strength of the sodium hydroxide solution in parts per hundred?

4. When perfectly dry sulphuric acid is treated with perfectly dry sodium hydroxide, no chemical change takes place. Explain.

5. When cold, concentrated sulphuric acid is added to zinc, no change takes place. Recall the action of dilute sulphuric acid on the same metal. How do you account for the difference?

6. A solution of hydrochloric acid in benzene does not conduct the electric current. When this solution is treated with zinc, will hydrogen be evolved? Explain.

7. (a) Write equation for preparation of hydrogen from zinc and dilute sulphuric acid. (b) Rewrite the same equation from the standpoint of the theory of electrolytic dissociation, (c) Subtract the common SO₄ ion from both members of the equation, (d) From the resulting equation, explain in what the preparation of hydrogen consists when examined from the standpoint of this theory.

8. In the same manner as in the preceding exercise, explain in what the action of sodium on water to give hydrogen consists.

[Pg 116]

CHAPTER XI

VALENCE

Definition of valence. A study of the formulas of various binary compounds shows that the elements differ between themselves in the number of atoms of other elements which they are able to hold in combination. This is illustrated in the formulas

It will be noticed that while one atom of chlorine combines with one atom of hydrogen, an atom of oxygen combines with two, an atom of nitrogen with three, one of carbon with four. The number which expresses this combining ratio between atoms is a definite property of each element and is called its valence.

DEFINITION: The valence of an element is that property which determines the number of the atoms of another element which its atom can hold in combination.

Valence a numerical property. Valence is therefore merely a numerical relation and does not convey any information in regard to the intensity of the affinity between atoms. Judging by the heat liberated in their union, oxygen has a far stronger affinity for hydrogen than does nitrogen, but an atom of oxygen can combine with two atoms only of hydrogen, while an atom of nitrogen can combine with three.[Pg 117]

Measure of valence. In expressing the valence of an element we must select some standard for comparison, just as in the measurement of any other numerical quantity. It has been found that an atom of hydrogen is never able to hold in combination more than one atom of any other element. Hydrogen is therefore taken as the standard, and other elements are compared with it in determining their valence. A number of other elements are like hydrogen in being able to combine with at most one atom of other elements, and such elements are called univalent. Among these are chlorine, iodine, and sodium. Elements such as oxygen, calcium, and zinc, which can combine with two atoms of hydrogen or other univalent elements, are said to be divalent. Similarly, we have trivalent, tetravalent, pentavalent elements. None have a valence of more than 8.

Indirect measure of valence. Many elements, especially among the metals, do not readily form compounds with hydrogen, and their valence is not easy to determine by direct comparison with the standard element. These elements, however, combine with other univalent elements, such as chlorine, and their valence can be determined from the compounds so formed.

Variable valence. Many elements are able to exert different valences under differing circumstances. Thus we have the compounds Cu₂O and CuO, CO and CO₂, FeCl₂ and FeCl₃. It is not always possible to assign a fixed valence to an element. Nevertheless each element tends to exert some normal valence, and the compounds in which it has a valence different from this are apt to be unstable and easily changed into compounds in which the valence of the element is normal. The valences of the various elements will become familiar as the elements are studied in detail.[Pg 118]

Valence and combining ratios. When elements combine to form compounds, the ratio in which they combine will be determined by their valences. In those compounds which consist of two elements directly combined, the union 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 above their symbols, we have the formulas

A divalent element, on the other hand, will combine with two atoms of a univalent element. Thus we have

(the numerals above each symbol representing the sum of the valences of the atoms of the element present). A trivalent atom will combine with three atoms of a univalent element, as in the compound

If a trivalent element combines with a divalent element, the union will be between two atoms of the trivalent element and three of the divalent element, since these numbers are the smallest which have equal valences. Thus the oxide of the trivalent metal aluminium has the formula Al₂O₃. Finally one atom of a tetravalent element such as carbon will combine with four atoms of a univalent element, as in the compound CH₄, or with two atoms of a divalent element, as in the compound CO₂.

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 replacing power of atoms. Just as elements having the same valence combine with each other atom for atom, so if they replace each other in a chemical reaction they will do so in the same ratio. This is seen in the following equations, in which a univalent hydrogen atom is replaced by a univalent sodium atom:

Similarly, one atom of divalent calcium will replace two atoms of univalent hydrogen or one of divalent zinc:

[Pg 120]

In like manner, one atom of a trivalent element will replace three of a univalent element, or two atoms will replace three atoms of a divalent element.

Valence and its applications to formulas of salts. While the true nature of valence is not understood and many questions connected with the subject remain unanswered, yet many of the main facts are of much help to the student. Thus the formula of a salt, differs from that of the acid from which it is derived in that the hydrogen of the acid has been replaced by a metal. If, then, it is known that a given metal forms a normal salt with a certain acid, the formula of the salt can at once be determined if the valence of the metal is known. Since sodium is univalent, the sodium salts of the acids HCl and H₂SO₄ will be respectively NaCl and Na₂SO₄. One atom of divalent zinc will replace 2 hydrogen atoms, so that the corresponding zinc salts will be ZnCl₂ and ZnSO₄.

The formula for aluminium sulphate is somewhat more difficult to determine. Aluminium is trivalent, and the simplest ratio in which the aluminium atom can replace the hydrogen in sulphuric acid is 2 atoms of aluminium (6 valences) to 3 molecules of sulphuric acid (6 hydrogen atoms). The formula of the sulphate will then be Al₂(SO₄)₃.

Valence and its application to equation writing. It will be readily seen that a knowledge of valence is also of very great assistance in writing the equations for reactions of double decomposition. Thus, in the general reaction between an acid and a base, the essential action is between the univalent hydrogen ion and the univalent hydroxyl ion. The base and the acid must always be taken in such proportions as to secure an equal number of each of these ions. Thus, in the reaction between ferric hydroxide (Fe(OH)₃) and sulphuric acid (H₂SO₄), it will be necessary to take 2 molecules of the former and 3 of the latter in order to have an equal number of the two ions, namely, 6. The equation will then be

Under certain conditions the salts Al₂(SO₄)₃ and CaCl₂ undergo double decomposition, the two metals, aluminium and calcium, exchanging places. The simplest ratio of exchange in this case is 2 atoms of aluminium (6 valences) and 3 atoms of calcium (6 valences).[Pg 121] The reaction will therefore take place between 1 molecule of Al₂(SO₄)₃ and 3 of CaCl₂, and the equation is as follows:

EXERCISES

1. Sodium, calcium, and aluminium have valences of 1, 2, and 3 respectively; write the formulas of their chlorides, sulphates, and phosphates (phosphoric acid = H₃PO₄), on the supposition that they form salts having the normal composition.

2. Iron forms one series of salts in which it has a valence of 2, and another series in which it has a valence of 3; write the formulas for the two chlorides of iron, also for the two sulphates, on the supposition that these have the normal composition.

3. Write the equation representing the neutralization of each of the following bases by each of the acids whose formulas are given:

4. Silver acts as a univalent element and calcium as a divalent element in the formation of their respective nitrates and chlorides. (a) Write the formula for silver nitrate; for calcium chloride. (b) When solutions of these two salts are mixed, the two metals, silver and calcium, exchange places; write the equation for the reaction.

5. Antimony acts as a trivalent element in the formation of a chloride. (a) What is the formula for antimony chloride? (b) When hydrosulphuric acid (H₂S) is passed into a solution of this chloride the hydrogen and antimony exchange places; write the equation for the reaction.

6. Lead has a valence of 2 and iron of 3 in the compounds known respectively as lead nitrate and ferric sulphate. (a) Write the formulas for these two compounds. (b) When their solutions are mixed the two metals exchange places; write the equation for the reaction.

[Pg 122]

CHAPTER XII

COMPOUNDS OF NITROGEN

Occurrence. As has been stated in a former chapter, nitrogen constitutes a large fraction of the atmosphere. The compounds of nitrogen, however, cannot readily be obtained from this source, since at any ordinary temperature nitrogen is able to combine directly with very few of the elements.

In certain forms of combination nitrogen occurs in the soil from which it is taken up by plants and built into complex substances composed chiefly of carbon, hydrogen, oxygen, and nitrogen. Animals feeding on these plants assimilate the nitrogenous matter, so that this element is an essential constituent of both plants and animals.

Decomposition of organic matter by bacteria. When living matter dies and undergoes decay complicated chemical reactions take place, one result of which is that the nitrogen of the organic matter is set free either as the element nitrogen, or in the form of simple compounds, such as ammonia (NH₃) or oxides of nitrogen. Experiment has shown that all such processes of decay are due to the action of different kinds of bacteria, each particular kind effecting a different change.

Decomposition of organic matter by heat. When organic matter is strongly heated decomposition into simpler substances takes place in much the same way as in the case of bacterial decomposition. Coal is a complex substance of[Pg 123] vegetable origin, consisting largely of carbon, but also containing hydrogen, oxygen, and nitrogen. When this is heated in a closed vessel so that air is excluded, about one seventh of the nitrogen is converted into ammonia, and this is the chief source from which ammonia and its compounds are obtained.

COMPOUNDS OF NITROGEN WITH HYDROGEN

Ammonia (NH₃). Several compounds consisting exclusively of nitrogen and hydrogen are known, but only one, ammonia, need be considered here.

Preparation of ammonia. Ammonia is prepared in the laboratory by a different method from the one which is used commercially.

1. Laboratory method. In the laboratory ammonia is prepared from ammonium chloride, a compound having the formula NH₄Cl, and obtained in the manufacture of coal gas. As will be shown later in the chapter, the group NH₄ in this compound acts as a univalent radical and is known as ammonium. When ammonium chloride is warmed with sodium hydroxide, the ammonium and sodium change places, the reaction being expressed in the following equation.

The ammonium hydroxide (NH₄OH) so formed is unstable and breaks down into water and ammonia.

Calcium hydroxide (Ca(OH)₂) is frequently used in place of the more expensive sodium hydroxide, the equations being

Fig. 35

2. Commercial method. Nearly all the ammonia of commerce comes from the gasworks. Ordinary illuminating gas is made by distilling coal, as will be explained later, and among the products of this distillation a solution of ammonia in water is obtained. This solution, known as gas liquor, contains not only ammonia but other soluble substances. Most of these combine chemically with lime, while ammonia does not; if then lime is added to the gas liquor and the liquor is heated, the ammonia is driven out from the mixture. It may be dissolved again in pure, cold water, forming aqua ammonia, or the ammonia water of commerce.

[Pg 125]

Physical properties. Under ordinary conditions ammonia is a gas whose density 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 purchased in liquid form in steel cylinders. The gas is colorless and has a strong, suffocating odor. It is extremely soluble in water, 1 l. of water at 0° and 760 mm. pressure dissolving 1148 l. of the gas. In dissolving this large volume of gas the water expands considerably, so that the density of the solution is less than that of water, the strongest solutions having a density of 0.88.

Chemical properties. Ammonia will not support combustion, nor will it burn under ordinary conditions. In an atmosphere of oxygen it burns with a feeble, yellowish flame. When quite dry it is not a very active substance, 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 by the application of pressure. If the pressure is removed from the liquid so obtained, it rapidly passes again into the gaseous state and in so doing absorbs a large amount of heat. Advantage is taken of this fact in the preparation of artificial ice. Large quantities of ammonia are also used 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[Pg 126] 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.

Fig. 36

Ammonium hydroxide (NH₄OH). The solution of ammonia in water is found to have strong basic properties and therefore contains hydroxyl ions. It turns red litmus blue; it has a soapy feel; it neutralizes acids, forming salts with them. It seems probable, therefore, that when ammonia dissolves in water it combines chemically with it according to the equation

and that it is the substance NH₄OH, called ammonium hydroxide, which has the basic properties, dissociating into the ions NH₄ and OH. Ammonium hydroxide has never been obtained in a pure state. At every attempt to isolate it the substance breaks up into water and ammonia,—

The ammonium radical. The radical NH₄ plays the part of a metal in many chemical reactions and is called ammonium. The ending -ium is given to the name to indicate the metallic properties of the substance, since the names[Pg 127] of the metals in general have that ending. The salts formed by the action of the base ammonium hydroxide on acids are called ammonium salts. Thus, with hydrochloric acid, ammonium chloride is formed in accordance with the equation

Similarly, with nitric acid, ammonium nitrate (NH₄NO₃) is formed, and with sulphuric acid, ammonium sulphate ((NH₄)₂S0₄).

It will be noticed that in the neutralization of ammonium hydroxide by acids the group NH₄ replaces one hydrogen atom of the acid, just as sodium does. The group therefore acts as a univalent metal.

Combination of nitrogen with hydrogen by volume. Under suitable conditions ammonia can be decomposed into nitrogen and hydrogen by passing electric sparks through the gas. Accurate measurement has shown that when ammonia is decomposed, two volumes of the gas yield one volume of nitrogen and three volumes of hydrogen. Consequently, if the two elements were to combine directly, one volume of nitrogen would combine with three volumes of hydrogen to form two volumes of ammonia. Here, as in the formation of steam from hydrogen and oxygen, small whole numbers serve to indicate the relation between the volumes of combining gases and that of the gaseous product.

COMPOUNDS OF NITROGEN WITH OXYGEN AND HYDROGEN

In addition to ammonium hydroxide, nitrogen forms several compounds with hydrogen and oxygen, of which nitric acid (HNO₃) and nitrous acid [Pg 128](HNO₂) are the most familiar.

Nitric acid (HNO₃). Nitric acid is not found to any extent in nature, but some of its salts, especially sodium nitrate (NaNO₃) and potassium nitrate (KNO₃) are found in large quantities. From these salts nitric acid can be obtained.

Fig. 37

Preparation of nitric acid. When sodium nitrate is treated with concentrated cold sulphuric acid, no chemical action seems to take place. If, however, the mixture is heated in a retort, nitric acid is given off as a vapor and may be easily condensed to a liquid by passing the vapor into a tube surrounded by cold water, as shown in Fig. 37. An examination of the liquid left in the retort shows that it contains sodium acid sulphate (NaHSO₄), so that the reaction may be represented by the equation

[Pg 129]

Fig. 38

Physical properties of nitric acid. Pure nitric acid is a colorless liquid, which boils at about 86° and has a density of 1.56. The concentrated acid of commerce contains about 68% of the acid, the remainder being water. Such a mixture has a density of 1.4. The concentrated acid fumes somewhat in moist air, and has a sharp choking odor.

Chemical properties. The most important chemical properties of nitric acid 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 blue litmus red and has a sour taste in dilute solutions. It forms hydrogen ions in solution and neutralizes bases forming salts. It also acts upon the oxides of most metals, forming a salt and water. It is one of the strongest acids.

2. Decomposition on heating. When boiled, or exposed for some time to sunlight, it suffers a partial decomposition according to the equation

The substance NO₂, called nitrogen peroxide, is a brownish gas, which is readily soluble in water and in nitric acid. It therefore dissolves in the undecomposed acid, and imparts a yellowish or reddish color to it. Concentrated[Pg 130] nitric acid highly charged with this substance is called fuming nitric acid.

3. Oxidizing action. According to its formula, nitric acid contains a large percentage of oxygen, and the reaction just mentioned shows that the compound is not a very stable one, easily undergoing decomposition. These properties should make it a good oxidizing agent, and we find that this is the case. Under ordinary circumstances, when acting as an oxidizing agent, it is decomposed according to the equation

The oxygen is taken up by the substance oxidized, and not set free, as is indicated in the equation. Thus, if carbon is oxidized by nitric acid, the oxygen combines with carbon, forming carbon dioxide (CO₂):

4. Action on metals. We have seen that when an acid acts upon a metal hydrogen is set free. Accordingly, when nitric acid acts upon a metal, such as copper, we should expect the reaction to take place which is expressed in the equation

This reaction does take place, but the hydrogen set free is immediately oxidized to water by another portion of the nitric acid according to the equation

As these two equations are written, two atoms of hydrogen are given off in the first equation, while three are used up in the second. In order that the hydrogen may be equal in[Pg 131] the two equations, we must multiply the first by 3 and the second by 2. We shall then have

The two equations may now be combined into one by adding the quantities on each side of the equality sign, canceling the hydrogen which is given off in the one reaction and used up in the other. We shall then have the equation

A number of other reactions may take place when nitric acid acts upon metals, resulting in the formation of other oxides of nitrogen, free nitrogen, or even ammonia. The reaction just given is, however, the usual one.

Salts of nitric acid,—nitrates. The salts of nitric acid are called nitrates. Many of these salts will be described in the study of the metals. They are all soluble in water, and when heated to a high temperature undergo decomposition. In a few cases a nitrate on being heated evolves oxygen, forming a nitrite:

[Pg 132]

In other cases the decomposition goes further, and the metal is left as oxide:

Nitrous acid (HNO₂). It is an easy matter to obtain sodium nitrite (NaNO₂), as the reaction given on the previous page indicates. Instead of merely heating the nitrate, it is better to heat it together with a mild reducing agent, such as lead, when the reaction takes place which is expressed by the equation

When sodium nitrite is treated with an acid, such as sulphuric acid, it is decomposed and nitrous acid is set free:

The acid is very unstable, however, and decomposes readily into water and nitrogen trioxide (N₂O₃):

Dilute solutions of the acid, however, can be obtained.

COMPOUNDS OF NITROGEN WITH OXYGEN

Nitrogen combines with oxygen to form five different oxides. The formulas and names of these are as follows:

These will now be briefly discussed.

Nitrous oxide (laughing gas) (N₂O). Ammonium nitrate, like all nitrates, undergoes decomposition when heated; and owing to the fact that it contains no metal, but does[Pg 133] contain both oxygen and hydrogen, the reaction is a peculiar one. It is represented by the equation

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 in water, and in solution has a slightly sweetish taste. It is easily converted into a liquid and can be purchased in this form. When inhaled it produces a kind of hysteria (hence the name "laughing gas"), and even unconsciousness and insensibility to pain if taken in large amounts. It has long been used as an anæsthetic for minor surgical operations, such as those of dentistry, but owing to its unpleasant after effects it is not so much in use now as formerly.

Chemically, nitrous oxide is remarkable for the fact that it is a very energetic oxidizing agent. Substances such as carbon, sulphur, iron, and phosphorus burn in it almost as brilliantly as in oxygen, forming oxides and setting free nitrogen. Evidently the oxygen in nitrous oxide cannot be held in very firm combination by the nitrogen.

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 takes place:

Nitric oxide is most conveniently prepared in this way. The metal is placed in the flask A (Fig. 39) and the acid added slowly through the funnel tube B. The gas escapes through C and is collected over water.[Pg 134]

Pure nitric oxide is a colorless gas, slightly heavier than air, and is practically insoluble in water. It is a difficult gas to liquefy. Unlike nitrous oxide, nitric oxide does not part with its oxygen easily, and burning substances introduced into this gas are usually extinguished. A few substances like phosphorus, which have a very strong affinity for oxygen and which are burning energetically in the air, will continue to burn in an atmosphere of nitric oxide. In this case the nitric oxide loses all of its oxygen and the nitrogen is set free as gas.

Action of nitric oxide with oxygen. When nitric oxide comes into contact with oxygen or with the air, it at once combines with the oxygen even at ordinary temperatures, forming a reddish-yellow gas of the formula NO₂, which is called nitrogen peroxide. This action is not energetic enough to produce a flame, though considerable heat is set free.

Nitrogen peroxide (NO₂). This gas, as we have just seen, is formed by allowing nitric oxide to come into contact with oxygen. It can also be made by heating certain nitrates, such as lead nitrate:

It is a reddish-yellow gas of unpleasant odor, which is quite poisonous when inhaled. It is heavier than air and is easily condensed to a liquid. It dissolves in water, but this solution is not a mere physical solution; the nitrogen peroxide is decomposed, forming a mixture of nitric and nitrous acids:

Nitrogen peroxide will not combine with more oxygen; it will, however, give up a part of its oxygen to burning substances, acting as an oxidizing agent:

[Pg 135]

Acid anhydrides. The oxides N₂O₃ (nitrogen trioxide) and N₂O₅ (nitrogen pentoxide) are rarely prepared and need not be separately described. They bear a very interesting relation to the acids of nitrogen. When dissolved in water they combine with the water, forming acids:

On the other hand, nitrous acid very easily decomposes, yielding water and nitrogen trioxide, and by suitable means nitric acid likewise may be decomposed into water and nitrogen pentoxide:

In view of the close relation between these oxides and the corresponding acids, they are called anhydrides of the acids, N₂O₃ being nitrous anhydride and N₂O₅ nitric anhydride.

DEFINITION: Any oxide which will combine with water to form an acid, or which together with water is formed by the decomposition of an acid, is called 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 concentrated sulphuric 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 than when strong bases are so neutralized. Suggest some possible cause for this.

4. Why is brine used in the manufacture of artificial ice?

5. Discuss the energy changes which take place in the manufacture of artificial ice.[Pg 136]

6. What weight of ammonium chloride is necessary to furnish enough ammonia to saturate 1 l. of water at 0° and 760 mm.?

7. What weight of sodium nitrate is necessary to prepare 100 cc. of commercial nitric acid? What weight of potassium nitrate is necessary to furnish the same weight of acid?

8. 100 l. of nitrogen peroxide were dissolved in water and neutralized with sodium hydroxide. What substances were formed and how much of each?(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 oxide under standard conditions?

11. (a) Calculate the percentage composition of the oxides of nitrogen. (b) What important law does this series of substances illustrate?

12. Write the equations representing the reactions between ammonium hydroxide, and sulphuric acid and nitric acid respectively, in accordance with the theory of electrolytic dissociation.

13. In the same way, write the equations representing the reactions between nitric acid and each of the following bases: NaOH, KOH, NH₄OH, Ca(OH)₂.

[Pg 137]

CHAPTER XIII

REVERSIBLE REACTIONS AND CHEMICAL EQUILIBRIUM

Reversible reactions. The reactions so far considered have been represented as continuing, when once started, until one or the other substance taking part in the reaction has been used up. In some reactions this is not the case. For example, we have seen that when steam is passed over hot iron the reaction is represented by the equation

On the other hand, when hydrogen is passed over hot iron oxide the reverse reaction takes place:

The reaction can therefore go in either direction, depending upon the conditions of the experiment. Such a reaction is called a reversible reaction. It is represented by an equation with double arrows in place of the equality sign, thus:

In a similar way, the equation

expresses the fact that under some conditions nitrogen may unite with hydrogen to form ammonia, while under other conditions ammonia decomposes into nitrogen and hydrogen.

The conversion of oxygen into ozone is also reversible and may be represented thus:

[Pg 138]

Chemical equilibrium. Reversible reactions do not usually go on to completion in one direction unless the conditions under which the reaction takes place are very carefully chosen. Thus, if iron and steam are confined in a heated tube, the steam acts upon the iron, producing iron oxide and hydrogen. But these substances in turn act upon each other to form iron and steam once more. When these two opposite reactions go on at such rates that the weight of the iron changed into iron oxide is just balanced by the weight of the iron oxide changed into iron, there will be no further change in the relative weights of the four substances present in the tube. The reaction is then said to have reached an equilibrium.

Factors which determine the point of equilibrium. There are two factors which have a great deal of influence in determining the point at which a given reaction will reach equilibrium.

1. Influence of the chemical nature of the substances. If two reversible reactions of the same general kind are selected, it has been found that the point of equilibrium is different in the two cases. For example, in the reactions represented by the equations

the equilibrium will be reached when very different quantities of the iron and zinc have been changed into oxides. The individual chemical properties of the iron and zinc have therefore marked influence upon the point at which equilibrium will be reached.

2. Influence of relative mass. If the tube in which the reaction

[Pg 139]

has come to an equilibrium is opened and more steam is admitted, an additional quantity of the iron will be changed into iron oxide. If more hydrogen is admitted, some of the oxide will be reduced to metal. The point of equilibrium is therefore dependent upon the relative masses of the substances taking part in the reaction. When one of the substances is a solid, however, its mass has little influence, since it is only the extent of its surface which can affect the reaction.

Conditions under which reversible reactions are complete. If, when the equilibrium between iron and steam has been reached, the tube is opened and a current of steam is passed in, the hydrogen is swept away as fast as it is formed. The opposing reaction of hydrogen upon iron oxide must therefore cease, and the action of steam on the iron will go on until all 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 oxide will be reduced to iron. A reversible reaction can therefore be completed in either direction when one of the products of the reaction is removed as fast as it is formed.

Equilibrium in solution. When reactions take place in solution in water the same general principles hold good. The matter is not so simple, however, as in the case just described, owing to the fact that many of the reactions in solution are due to the presence of ions. The substances most commonly employed in solution are acids, bases, or salts, and all of these undergo dissociation. Any equilibrium which may be reached in solutions of these substances must take place between the various ions formed, on the[Pg 140] one hand, and the undissociated molecules, on the other. Thus, when nitric acid is dissolved in water, equilibrium is reached in accordance with the equation

Conditions under which reversible reactions in solution are complete. The equilibrium between substances in solution may be disturbed and the reaction caused to go on in one direction to completion in either of three ways.

1. A gas may be formed which escapes from the solution. When sodium nitrate and sulphuric acid are brought together in solution all four ions, Na⁺, NO₃^-, H⁺, SO₄^-, are formed. These ions are free to rearrange themselves in various combinations. For example, the H⁺ and the NO₃^- ions will reach the equilibrium

If the experiment is performed with very little water present, as is the case in the preparation of nitric acid, the equilibrium will be reached when most of the H⁺ and the NO₃^- ions have combined to form undissociated HNO₃.

Finally, if the mixture is now heated above the boiling point of nitric acid, the acid distills away as fast as it is formed. More and more H⁺ and NO₃^- ions will then combine, and the process will continue until one or the other of them has all been removed from the solution. The substance remaining is sodium acid sulphate (NaHSO₄), and the reaction can therefore be expressed by the equation

2. An insoluble solid may be formed. When hydrochloric acid (HCl) and [Pg 141]silver nitrate (AgNO₃) are brought together in solution the following ions will be present: H⁺, Cl^-, Ag⁺, NO₃^-. The ions Ag⁺ and Cl^- will then set up the equilibrium

But silver chloride (AgCl) is almost completely insoluble in water, and as soon as a very little of it has formed the solution becomes supersaturated, and the excess of the salt precipitates. More silver and chlorine ions then unite, and this continues until practically all of the silver or the chlorine ions have been removed from the solution. We then say that the following reaction is complete:

3. Two different ions may form undissociated molecules. In the neutralization of sodium hydroxide by hydrochloric acid the ions H⁺ and OH^- come to the equilibrium

But since water is almost entirely undissociated, equilibrium can only be reached when there are very few hydroxyl or hydrogen ions present. Consequently the two ions keep uniting until one or the other of them is practically removed from the solution. When this occurs the neutralization expressed in the following equation is complete:

Preparation of acids. The principle of reversible reactions finds practical application in the preparation of most of the common acids. An acid is usually prepared by treating the most common of its salts with some other acid of high boiling point. The mixture is then heated until the lower boiling acid desired distills out. Owing to[Pg 142] its high boiling point (338°), sulphuric acid is usually employed for this purpose, most other acids boiling below that temperature.

EXERCISES

1. What would take place when solutions of silver nitrate and sodium chloride are brought together? What other chlorides would act in the same way?

2. Is the reaction expressed by the equation NH₃ + H₂O = NH₄OH reversible? If so, state the conditions under which it will go in each direction.

3. Is the reaction expressed by the equation 2H + O = H₂O reversible? If so, state the conditions under which it will go in each direction.

4. Suggest a method for the preparation of hydrochloric acid.

[Pg 143]

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 quantities in the uncombined form, especially in the neighborhood of volcanoes. Sicily has long been famous for its sulphur mines, and smaller deposits are found in Italy, Iceland, Mexico, and especially in Louisiana, where it is mined extensively. In combination, sulphur occurs abundantly in the form of sulphides and sulphates. In smaller amounts it is found in a great variety of minerals, and it is a constituent of many animal and vegetable substances.

Extraction of sulphur. Sulphur is prepared from the native substance, the separation of crude sulphur from the rock and earthy materials with which it is mixed being a very simple process. The ore from the mines is merely heated until the sulphur melts and drains away from the earthy impurities. The crude sulphur obtained in this way is distilled in a retort-shaped vessel made of iron, the exit tube of which opens into a cooling chamber of brickwork. When the sulphur vapor first enters the cooling chamber it condenses as a fine crystalline powder called flowers of sulphur. As the condensing chamber becomes warm, the sulphur collects as a liquid in it, and is drawn off into cylindrical molds, the product being called roll sulphur or brimstone.[Pg 144]

Physical properties. Roll sulphur is a pale yellow, crystalline solid, without marked taste and with but a faint odor. It is insoluble in water, but is freely soluble in a few liquids, notably in carbon disulphide. Roll sulphur melts at 114.8°. Just above the melting point it forms a rather thin, straw-colored liquid. As the temperature is raised, this liquid turns darker in color and becomes thicker, until at about 235° it is almost black and is so thick that the vessel containing it can be inverted without danger of the liquid running out. At higher temperatures it becomes thin once more, and boils at 448°, forming a yellowish vapor. On cooling the same changes take place in reverse order.

Varieties of sulphur. Sulphur is known in two general forms, crystalline and 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 in carbon disulphide it separates in crystals which have the same color and melting point as roll sulphur, and are rhombic in shape. Roll sulphur is made up of minute rhombic crystals.

2. Monoclinic sulphur. When melted sulphur is allowed to cool until a part of the liquid has solidified, and the remaining liquid is then poured off, it is found that the solid sulphur remaining in the vessel has assumed the form of fine needle-shaped crystals. These differ much in appearance from the rhombic crystals obtained by crystallizing sulphur from its solution in carbon disulphide. The needle-shaped form is called monoclinic sulphur. The two varieties differ also in density and in melting point, the monoclinic sulphur melting at 120°.[Pg 145]

Monoclinic and rhombic sulphur remain unchanged in contact with each other at 96°. Above this temperature the rhombic changes into monoclinic; 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 readily obtained. These are white sulphur and plastic sulphur.

1. White sulphur. Flowers of sulphur, the preparation of which has been described, consists of a mixture of rhombic crystals and amorphous particles. 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 it assumes a gummy, doughlike form, which is quite elastic. This can be seen in a very striking manner by distilling sulphur from a small, short-necked retort, such as is represented in Fig. 40, and allowing the liquid to run directly into water. In a few days it becomes quite brittle and passes over into ordinary rhombic sulphur.

Fig. 40

Chemical properties of sulphur. When sulphur is heated to its kindling temperature in oxygen or in the air it burns with a pale blue flame, forming sulphur dioxide (SO₂). Small quantities of sulphur trioxide [Pg 146](SO₃) may also be formed in the combustion of sulphur. Most metals when heated with sulphur combine directly with it, forming metallic sulphides. In some cases the action is so energetic that the mass becomes incandescent, as has been seen in the case of iron uniting with sulphur. This property recalls the action of oxygen upon metals, and in general the metals which combine readily with oxygen are apt to combine quite readily with sulphur.

Uses of sulphur. Large quantities of sulphur are used as a germicide in vineyards, also in the manufacture of gunpowder, matches, vulcanized rubber, and sulphuric acid.

COMPOUNDS OF SULPHUR WITH HYDROGEN

Hydrosulphuric acid (H₂S). This substance is a gas having the composition expressed by the formula H₂S and is commonly called hydrogen sulphide. It is found in the vapors issuing from volcanoes, and in solution in the so-called sulphur waters of many springs. It is formed when organic matter containing sulphur undergoes decay, just as ammonia is formed under similar circumstances from nitrogenous matter.

Preparation. Hydrosulphuric acid is prepared in the laboratory by treating a sulphide with an acid. Iron sulphide (FeS) is usually employed:

A convenient apparatus is shown in Fig. 41. A few lumps of iron sulphide are placed in the bottle A, and dilute acid is added in small quantities at a time through the funnel tube B, the gas escaping through the tube C.

Fig. 41

[Pg 147]

Physical properties. Hydrosulphuric acid is a colorless gas, having a weak, disagreeable taste and an exceedingly offensive odor. It is rather sparingly soluble in water at ordinary temperatures, about three volumes dissolving 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 largely with air produces headache, dizziness, and nausea. It is a little heavier than air, having a density of 1.18.

Chemical properties. The most important chemical properties of hydrosulphuric acid are the following:

1. Acid properties. Hydrosulphuric acid is a weak acid. In solution in water it turns blue litmus red and neutralizes bases, forming salts called sulphides.

2. Action on oxygen. The elements composing hydrosulphuric acid have each a strong affinity for oxygen, and are not held together very firmly. Consequently the gas burns readily in oxygen or the air, according to the equation

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:

[Pg 148]

3. Reducing action. Owing to the ease with which hydrosulphuric acid decomposes and the strong affinity of both sulphur and hydrogen for oxygen, the substance is a strong reducing agent, taking oxygen away from many substances which contain it.

4. Action on metals. Hydrosulphuric acid acts towards metals in a way very similar to water. Thus, when it is passed over heated iron in a tube, the reaction is represented by the equation

Water in the form of steam, under similar circumstances, acts according to the equation

Salts of hydrosulphuric acid,—sulphides. The salts of hydrosulphuric acid, called sulphides, form an important class of salts. Many of them are 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 are insoluble in acids. Consequently, when hydrosulphuric acid is passed into a solution of a salt, it often happens that a sulphide is precipitated. With copper chloride the equation is

Because of the fact that some metals are precipitated in this way as sulphides while others are not, hydrosulphuric acid is extensively used in the separation of the metals in the laboratory.

[Pg 149]

OXIDES OF SULPHUR

Sulphur forms two well-known compounds with oxygen: sulphur dioxide (SO₂), sometimes called sulphurous anhydride; and sulphur trioxide (SO₃), frequently called sulphuric anhydride.

Sulphur dioxide (SO₂). Sulphur dioxide occurs in nature in the gases issuing from volcanoes, and in solution in the water of many springs. It is likely to be found wherever sulphur compounds are undergoing oxidation.

Preparation. Three general ways may be mentioned for the preparation of sulphur dioxide:

1. By the combustion of sulphur. Sulphur dioxide is readily formed by the combustion of sulphur in oxygen or the air:

It is also formed when substances containing sulphur are burned:

2. By the reduction of sulphuric acid. When concentrated sulphuric acid is heated with certain metals, such as copper, part of the acid is changed into copper sulphate, and part is reduced to sulphurous acid. The latter then decomposes into sulphur dioxide and water, the complete equation being

[Pg 150]

3. By the action of an acid on a sulphite. Sulphites are salts of sulphurous acid (H₂SO₃). When a sulphite is treated with an acid, sulphurous acid is set free, and being very unstable, decomposes into water and sulphur dioxide. These reactions are expressed in the equations

Explanation of the reaction. In this case we have two reversible reactions depending on each other. In the first reaction,

we should expect an equilibrium to result, for none of the four substances in the equation are insoluble or volatile when water is present to hold them in solution. But the quantity of the H₂SO₃ is constantly diminishing, owing to the fact that it decomposes, as represented in the equation

and the sulphur dioxide, being a gas, escapes. No equilibrium can therefore result, since the quantity of the sulphurous acid is constantly being diminished because of the escape of sulphur dioxide.

Physical properties. Sulphur dioxide is a colorless gas, which at ordinary 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 water dissolving eighty of the gas under standard conditions. It is easily condensed to a colorless liquid, and can be purchased in this condition stored in strong bottles, such as the one represented in Fig. 42.

Fig. 42

Chemical properties. Sulphur dioxide has a marked tendency to combine with other substances, and is therefore an[Pg 151] active substance chemically. It combines with oxygen gas, but not very easily. It can, however, take oxygen away from some other substances, and is therefore a good reducing agent. Its most marked chemical property is its ability to combine with water to form sulphurous acid (H₂SO₃).

Sulphurous acid (H₂SO₃). When sulphur dioxide dissolves in water it combines chemically with it to form sulphurous acid, an unstable substance having the formula H₃SO₃. It is impossible to prepare this acid in pure form, as it breaks down very easily into water and sulphur dioxide. The reaction is therefore reversible, and is expressed by the equation

Solutions of the acid in water have a number of interesting properties.

1. Acid properties. The solution has all the properties typical of an acid. When neutralized by bases, sulphurous acid yields a series of salts called sulphites.

2. Reducing properties. Solutions of sulphurous acid act as good reducing agents. This is due to the fact that sulphurous acid has the power of taking up oxygen from the air, or from substances rich in oxygen, and is changed by this reaction into sulphuric acid:

3. Bleaching properties. Sulphurous acid has strong bleaching properties, acting upon many colored substances in such a way as to destroy their color. It is on this account used to bleach paper, straw goods, and even such foods as canned corn.

4. Antiseptic properties. Sulphurous acid has marked antiseptic properties, and on this account has the power[Pg 152] of arresting fermentation. It is therefore used as a preservative.

Salts of sulphurous acid,—sulphites. The sulphites, like sulphurous acid, have the power of taking up oxygen very readily, and are good reducing agents. On account of this tendency, commercial sulphites are often contaminated with sulphates. A great deal of sodium sulphite is used in the bleaching industry, and as a reagent for softening paper pulp.

Sulphur trioxide (SO₃). When sulphur dioxide and oxygen are heated together at a rather high temperature, a small amount of sulphur trioxide (SO₃) is formed, but the reaction is slow and incomplete. If, however, the heating takes place in the presence of very fine platinum dust, the reaction is rapid and nearly complete.

Fig. 43

Properties of sulphur trioxide. Sulphur trioxide is a colorless liquid, which solidifies at about 15° and boils at 46°.[Pg 153] A trace of moisture causes it to solidify into a mass of silky white crystals, somewhat resembling asbestos fiber in appearance. In contact with the air it fumes strongly, and when thrown upon water it dissolves with a hissing sound and the liberation of a great deal of heat. The product of this reaction is sulphuric acid, so that sulphur trioxide is the anhydride of that acid:

Catalysis. It has been found that many chemical reactions, such as the union of sulphur dioxide with oxygen, are much influenced by the presence of substances which do not themselves seem to take a part in the reaction, and are left apparently unchanged after it has ceased. These reactions go on very slowly under ordinary circumstances, but are greatly hastened by the presence of the foreign substance. Substances which hasten very slow reactions in this way are said to act as catalytic 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 a reaction, but does not change its products.

Examples of Catalysis. We have already had several instances of such action. Oxygen and hydrogen combine with each other at ordinary temperatures in the presence of platinum powder, while if no catalytic agent is present they do not combine in appreciable quantities until a rather high temperature is reached. Potassium chlorate, when heated with manganese dioxide, gives up its oxygen at a much lower temperature than when heated alone. Hydrogen dioxide decomposes very rapidly when powdered manganese dioxide is sifted into its concentrated solution.[Pg 154]

On the other hand, the catalytic agent sometimes retards chemical action. For example, a solution of hydrogen dioxide decomposes more slowly when it contains a little phosphoric acid than when perfectly pure. For this reason commercial hydrogen dioxide always contains phosphoric acid.

Many reactions are brought about by the catalytic action of traces of water. For example, phosphorus will not burn in oxygen in the absence of all moisture. Hydrochloric acid will not unite with ammonia if the reagents are perfectly dry. It is probable that many of the chemical transformations in physiological processes, such as digestion, are assisted by certain substances acting as catalytic agents. The principle of catalysis is therefore very important.

Sulphuric acid (oil of vitriol) (H₂SO₄). Sulphuric acid is one of the most important of all manufactured chemicals. Not only is it one of the most common reagents in the laboratory, but enormous quantities of it are used in many of the industries, especially in the refining of petroleum, the manufacture of nitroglycerin, sodium carbonate, and fertilizers.

Manufacture of sulphuric acid. 1. Contact process. The reactions taking place in this process are represented by the following equations:

To bring about the first of these reactions rapidly, a catalyzer is employed, and the process is carried out in the following way: Large iron tubes are packed with some porous material, such as calcium and magnesium sulphates, which contains a suitable catalytic substance scattered through it. The catalyzers most used are platinum powder,[Pg 155] vanadium oxide, and iron oxide. Purified sulphur dioxide and air are passed through the tubes, which are kept at a temperature of about 350°. Sulphur trioxide is formed, and as it issues from the tube it is absorbed in water or dilute sulphuric acid. The process is continued until all the water in the absorbing vessel has been changed into sulphuric acid, so that a very concentrated acid is made in this way. An excess 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 employed until recent years, and still in very extensive use, is much more complicated. The reactions are quite involved, but the conversion of water, sulphur dioxide, and oxygen into sulphuric acid is accomplished by the catalytic action of oxides of nitrogen. The reactions are brought about in large lead-lined chambers, into which oxides of nitrogen, sulphur dioxide, steam, and air are introduced in suitable proportions.

[Pg 156]

The relation between sulphuric acid and nitrosulphuric acid can be seen by comparing their structural formulas:

The latter may be regarded as derived from the former by the substitution of the nitro group (NO₂) for the hydroxyl group (OH).

Fig. 44.

The sulphuric acid plant. Fig. 44 illustrates the simpler parts of a plant used in the manufacture of sulphuric acid by the chamber process. Sulphur or some sulphide, as FeS₂, is burned in furnace A. The resulting sulphur dioxide, together with air and some nitrogen peroxide, are conducted into the large chambers, the capacity of each chamber being about 75,000 cu. ft. Steam is also admitted into these chambers at different points. These compounds react to form sulphuric acid, according to the equations given above. The nitrogen left after the withdrawal of the oxygen from the admitted air escapes through the Gay-Lussac tower X. In order to prevent the escape of the oxides of nitrogen regenerated in the reaction, the tower is filled with lumps of coke, over which trickles concentrated sulphuric acid admitted from Y. The nitrogen peroxide dissolves in the acid and the resulting solution collects in H. This is pumped into E, where it is mixed with dilute acid and allowed to trickle down through the chamber D (Glover tower), which is filled with some acid-resisting rock. Here the nitrogen peroxide is expelled from the solution by the action of the hot gases entering from A, and together with them enters the first chamber again. The acid from which the nitrogen peroxide is expelled collects in F. Theoretically, a small amount of nitrogen peroxide would suffice to prepare an unlimited amount of sulphuric acid; practically, some of it escapes, and this is replaced by small amounts admitted at B.[Pg 157]

The sulphuric acid so formed, together with the excess of condensed steam, collect upon the floor of the chambers in the form of a liquid containing from 62% to 70% of sulphuric acid. The product is called chamber acid and is quite impure; but for many purposes, such as the manufacture of fertilizers, it needs no further treatment. It can be concentrated by boiling it in vessels made of iron or platinum, which resist the action of the acid, nearly all the water boiling off. Pure concentrated acid can be made best by the contact process, while the chamber process is cheaper for the dilute impure acid.

Physical properties. Sulphuric acid is a colorless, oily liquid, nearly twice as heavy as water. The ordinary concentrated acid contains about 2% of water, has a density of 1.84, and boils at 338°. It is sometimes called oil of vitriol, since it was formerly made by distilling a substance called green vitriol.

Chemical properties. Sulphuric acid possesses chemical properties which make it one of the most important of chemical substances.

1. Action as an acid. In dilute solution sulphuric acid acts as any other acid, forming salts with oxides and hydroxides.

2. Action as an oxidizing agent. Sulphuric acid contains a large percentage of oxygen and is, like nitric acid, a very good oxidizing agent. When the concentrated acid is heated with sulphur, carbon, and many other substances, oxidation takes place, the sulphuric acid decomposing according to the equation

3. Action on metals. In dilute solution sulphuric acid acts upon many metals, such as zinc, forming a sulphate and liberating hydrogen. When the concentrated acid is employed the hydrogen set free is oxidized by a new portion[Pg 158] of the acid, with the liberation of sulphur dioxide. With copper the reactions are expressed by the equations

By combining these equations the following one is obtained:

4. Action on salts. We have repeatedly seen that an acid of high boiling point heated with the salt of some acid of lower boiling point will 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 used largely in the preparation of other acids.

5. Action on water. Concentrated sulphuric acid has a very great affinity for water, and is therefore an effective dehydrating agent. Gases which have no chemical action upon sulphuric acid can be freed from water vapor by bubbling them through the strong acid. When the acid is diluted with water much heat is set free, and care must be taken to keep the liquid thoroughly stirred during the mixing, and to pour the acid into the water,—never the reverse.

Not only can sulphuric acid absorb water, but it will often withdraw the elements hydrogen and oxygen from a compound containing them, decomposing the compound, and combining with the water so formed. For this reason most organic substances, such as sugar, wood, cotton, and woolen fiber, and even flesh, all of which contain much oxygen and hydrogen in addition to carbon, are charred or burned by the action of the concentrated acid.[Pg 159]

Salts of sulphuric acid,—sulphates. The sulphates form a very important class of salts, and many of them have commercial uses. Copperas (iron sulphate), blue vitriol (copper sulphate), and Epsom salt (magnesium sulphate) serve as examples. Many sulphates are important minerals, prominent among these being gypsum (calcium sulphate) and barytes (barium sulphate).

Monobasic and dibasic acids. Such acids as hydrochloric and nitric acids, which have only one replaceable hydrogen atom in the molecule, or in other words yield one hydrogen ion in solution, are called monobasic acids. Acids yielding two hydrogen ions in solution are called dibasic acids. Similarly, we may have tribasic and tetrabasic acids. The three acids of sulphur are dibasic acids. It is therefore possible for each of them to form both normal and acid salts. The acid salts can be made in two ways: the acid may be treated with only half enough base to neutralize it,—

[Pg 160]

or a normal salt may be treated with the free acid,—

Acid sulphites and sulphides may be made in the same ways.

Carbon disulphide (CS₂). When sulphur vapor is passed over highly heated carbon the two elements combine, forming carbon disulphide (CS₂), just as oxygen and carbon unite to form carbon dioxide (CO₂). The substance is a heavy, colorless liquid, possessing, when pure, a pleasant ethereal odor. On standing for some time, especially when exposed to sunlight, it undergoes a slight decomposition and acquires a most disagreeable, rancid odor. It has the property of dissolving many substances, such as gums, resins, and waxes, which are insoluble in most liquids, and it is extensively used as a solvent for such substances. It is also used as an insecticide. It boils at a low temperature (46°), and its vapor is very inflammable, burning in the air to form carbon dioxide and sulphur dioxide, according to the equation

Fig. 45

[Pg 161]

Comparison of sulphur and oxygen. A comparison of the formulas and the chemical properties of corresponding compounds of oxygen and sulphur brings to light many striking similarities. The conduct of hydrosulphuric acid and water toward many substances has been seen to be very similar; the oxides and sulphides of the metals have analogous formulas and undergo many parallel reactions. Carbon dioxide and disulphide are prepared in similar ways and undergo many analogous reactions. It is clear, therefore, that these two elements are far more closely related to each other than to any of the other elements so far studied.

Selenium and tellurium. These two very uncommon elements are still more closely related to sulphur than is oxygen. They occur in comparatively small quantities and are usually found associated with sulphur and sulphides, either as the free elements or more commonly in combination with metals. They form compounds with hydrogen of the formulas H₂Se and H₂Te; these bodies are gases with properties very similar to those of H₂S. They also form oxides and oxygen acids which resemble the corresponding sulphur compounds. The elements even have allotropic forms corresponding very closely to those of sulphur. Tellurium is sometimes found in combination with gold and copper, and occasions some difficulties in the refining of these metals. The elements have very few practical applications.

Crystallography. In order to understand the difference between the two kinds of sulphur crystals, it is necessary to know something about crystals in general and the forms which they may assume. An examination of a large number of crystals has shown that although they may differ much in geometric form, they can all be considered as modifications of a few simple plans. The best way to understand the relation of one crystal to another is to look upon every crystal as having its faces and angles arranged in definite fashion about[Pg 162] certain imaginary lines drawn through the crystal. These lines are called axes, and bear much the same relation to a crystal as do the axis and parallels of latitude and longitude to the earth and a geographical study of it. All crystals can be referred to one of six simple plans or systems, which have their axes as 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 right angles.

Fig. 46

2. Tetragonal system (Fig. 47). Two equal axes and one of different length, all at right angles to each other.

Fig. 47

3. Orthorhombic system (Fig. 48). Three unequal axes, all at right angles to each other.

Fig. 48

4. Monoclinic system (Fig. 49). Two axes at right angles, and a third at right angles to one of these, but inclined to the other.

Fig. 49

5. Triclinic system (Fig. 50). Three axes, all inclined to each other.

Fig. 50

6. Hexagonal system (Fig. 51). Three equal axes in the same plane intersecting at angles of 60°, and a fourth at right angles to all of these.

Fig. 51

Every crystal can be imagined to have its faces and angles arranged in a definite 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; while in a regular octohedron an axis ends in each solid angle. These forms are shown in Fig. 46. It will be seen that both of these figures belong to the same system, though they are very different in appearance. In the same way, many geometric[Pg 163] forms may be derived from each of the systems, and the light lines about the axes in the drawings show two of the simplest forms of each of the systems.

In general a given substance always crystallizes in the same system, and two corresponding faces of each crystal of it always make the same angle with each other. A few substances, of which sulphur is an example, crystallize in two different systems, and the crystals differ in such physical properties as melting point and density. Such substances are said to be dimorphous.

EXERCISES

1. (a) Would the same amount of heat be generated by the combustion of 1 g. of each of the allotropic modifications of sulphur? (b) Would the same amount of sulphur dioxide be formed in each case?

2. Is the equation for the preparation of hydrosulphuric acid a reversible one? As ordinarily carried out, does the reaction complete itself?

3. Suppose that hydrosulphuric acid were a liquid, would it be necessary to modify the method of preparation?

4. Can sulphuric acid be used to dry hydrosulphuric acid? Give reason for answer.

5. Does dry hydrosulphuric acid react with litmus paper? State reason for answer.

6. How many grams of iron sulphide are necessary to prepare 100 l. of hydrosulphuric acid when the laboratory conditions are 17° and 740 mm. pressure?

7. Suppose that the hydrogen in 1 l. of hydrosulphuric acid were liberated; what volume would it occupy, the gases being measured under the same conditions?

8. Write the equations representing the reaction between hydrosulphuric acid and sodium hydroxide and ammonium hydroxide respectively.

9. Show that the preparation of sulphur dioxide from a sulphite is similar in principle to the preparation of hydrogen sulphide.

10. (a) Does dry sulphur dioxide react with litmus paper? (b) How can it be shown that a solution of sulphur dioxide in water acts like an acid?[Pg 164]

11. (a) Calculate the percentage composition of sulphurous anhydride and sulphuric anhydride. (b) Show how these two substances are in harmony with the law of multiple proportion.

12. How many pounds of sulphur would be necessary in the preparation of 100 lb. of 98% sulphuric acid?

13. What weight of sulphur dioxide is necessary in the preparation of 1 kg. of sodium sulphite?

14. What weight of copper sulphate crystals can be obtained by dissolving 1 kg. of copper in sulphuric acid and crystallizing the product from water?

15. Write the names and formulas of the oxides and oxygen acids of selenium and tellurium.

16. In the commercial preparation of carbon disulphide, what is the function of the electric current?

17. If the Gay-Lussac tower were omitted from the sulphuric acid factory, what effect would this have on the cost of production of sulphuric acid?

[Pg 165]

CHAPTER XV

PERIODIC LAW

A number of the elements have now been studied somewhat closely. The first three of these, oxygen, hydrogen, and nitrogen, while having some physical properties in common with each other, have almost no point of similarity as regards their chemical conduct. On the other hand, oxygen and sulphur, while quite different physically, have much in common in their chemical properties.

About eighty elements are now known. If all of these should have properties as diverse as do oxygen, hydrogen, and nitrogen, the study of chemistry would plainly be a very difficult and complicated one. If, however, the elements can be classified in groups, the members of which have very similar properties, the study will be very much simplified.

Earlier classification of the elements. Even at an early period efforts were made to discover some natural principle in accordance with which the elements could be classified. Two of these classifications may be mentioned here.

1. Classification into metals and non-metals. The classification into metals and non-metals most naturally suggested itself. This grouping was based 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 in character, and the ability to form bases came to be regarded as a characteristic property of the metals. The[Pg 166] non-metals possessed physical properties which were the reverse of those of the metals, and were 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 as iodine, have a high luster; some elements can form either an acid or a base.

2. Classification into triad families. In 1825 Döbereiner observed that an interesting relation exists between the atomic weights of chemically similar elements. To illustrate, lithium, sodium, and potassium resemble each other very closely, and the atomic weight of sodium is almost exactly an arithmetical mean between those of the other two: (7.03 + 39.15)/2 = 23.09. In many chemical and physical properties sodium is midway between the other two.

A number of triad families were found, but among eighty elements, whose atomic weights range all the way from 1 to 240, such agreements might be mere chance. Moreover many elements did not appear to belong to such families.

Periodic division. In 1869 the Russian chemist Mendeléeff devised an arrangement of the elements based on their atomic weights, which has proved to be of great service in the comparative study of the elements. A few months later the German, Lothar Meyer, independently suggested the same ideas. This arrangement brought to light a great generalization, now known as the periodic law. An exact statement of the law will be given after the method of arranging the elements has been described.

DMITRI IVANOVITCH MENDELÉEFF (Russian) (1834-1907)

Author of the periodic law; made many investigations on the physical constants of elements and compounds; wrote an important book entitled "Principles of Chemistry"; university professor and government official

[Pg 167]

Arrangement of the periodic table. The arrangement suggested by Mendeléeff, modified somewhat by more recent investigations, is as follows: Beginning with lithium, which has an atomic weight of 7, the elements are arranged in a horizontal row in the order of their atomic weights, 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 eighth element, sodium, is very similar to lithium. It is placed just under lithium, 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 under sodium, 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, and potassium, which have been placed in a vertical row because of this resemblance, but the elements in the other vertical rows exhibit much of the same kind of similarity among themselves, and evidently form little natural groups.

The three elements following manganese, namely, iron, nickel, and cobalt, have atomic weights near together, and are very similar chemically. They do not strongly resemble any of the elements so far considered, and are accordingly placed in a group by themselves, following manganese. A new row is begun with copper, which somewhat resembles the elements of the first vertical column. Following the fifth and seventh rows are groups of three closely related elements, so that the completed arrangement has the appearance represented in the table on page 168.

[Pg 168]

THE PERIODIC ARRANGEMENT OF THE ELEMENTS

[Pg 169]

Place of the atmospheric elements. When argon was discovered it was seen at once that there was no place in the table for an element of atomic weight approximately 40. When the other inactive elements were found, however, it became apparent that they form a group just preceding Group 1. They are accordingly arranged in this way in Group 0 (see table on opposite page). A study of this table brings to light certain very striking facts.

Properties of elements vary with atomic weights. There is evidently a close relation between the properties of an element and its atomic weight. Lithium, at the beginning of the first group, is a very strong base-forming element, with pronounced metallic properties. Beryllium, following lithium, is less strongly base-forming, while boron has some base-forming and some acid-forming properties. In carbon all base-forming properties have disappeared, and the acid-forming properties are more marked than in boron. These become still more emphasized as we pass through nitrogen and oxygen, until on reaching fluorine we have one of the strongest acid-forming elements. The properties of these seven elements therefore vary regularly with their atomic weights, or, in mathematical language, are regular functions of them.

Periodic law. The properties of the first seven elements vary continuously—that is steadily—away from base-forming and toward acid-forming properties. If lithium had the smallest atomic weight of any of the elements, and fluorine the greatest, so that in passing from one to the other we had included all the elements, we could say that the properties of elements are continuous functions of their atomic weights. But fluorine is an element of small atomic weight, and the one following it, sodium, breaks the regular order, for in it reappear all the characteristic properties of lithium. Magnesium, following sodium, bears much the same relation to[Pg 170] beryllium that sodium does to lithium, and the properties of the elements in the second row vary much as they do in the first row until potassium is reached, when another repetition begins. The properties of the elements do not vary continuously, therefore, with atomic weights, but at regular intervals there is a repetition, or period. This generalization is known as the periodic law, and may be stated thus: The properties of elements are periodic functions of their atomic weights.

The two families in a group. While all the elements in a given vertical column bear a general resemblance to each other, it has been noticed that those belonging to periods having even numbers are very strikingly similar to each other. They are placed at the left side of the group columns. In like manner, the elements belonging to the odd periods are very similar and are arranged at the right side of the group columns. Thus calcium, strontium, and barium are very much alike; so, too, are magnesium, zinc, and cadmium. The resemblance between calcium and magnesium, or strontium and zinc, is much less marked. This method of arrangement therefore divides each group into two families, each containing four or five members, between which there is a great similarity.

Family resemblances. Let us now inquire more closely in what respects the elements of a family resemble each other.

1. Valence. In general the valence of the elements in a family is the same, and the formulas of their compounds are therefore similar. If we know that the formula of sodium chloride is NaCl, it is pretty certain that the formula of potassium chloride will be KCl—not KCl₂ or [Pg 171]KCl₃. The general formulas R₂O, RO, etc., placed below the columns show the formulas of the oxides of the elements in the column provided they form oxides. In like manner the formulas RH, RH₂, etc., show the composition of the compounds formed with hydrogen or chlorine.

2. Chemical properties. The chemical properties of the members of a family are quite similar. If one member is a metal, the others usually are; if one is a non-metal, so, too, are the others. The families in the first two columns consist of metals, while the elements found in the last two columns form acids. There is in addition a certain regularity in properties of the elements in each family. If the element at the head of the family is a strong acid-forming element, this property is likely to diminish gradually, as we pass to the members of the family with higher atomic weights. Thus phosphorus is strongly acid-forming, arsenic less so, antimony still less so, while bismuth has almost no acid-forming properties. We shall meet with many illustrations of this fact.

3. Physical properties. In the same way, the physical properties of the members of a family are in general somewhat similar, and show a regular gradation as we pass from element to element in the family. Thus the densities of the members of the magnesium family are

Their melting points are

Value of the periodic law. The periodic law has proved of much value in the development of the science of chemistry.

1. It simplifies study. It is at once evident that such regularities very much simplify the study of chemistry.[Pg 172] A thorough study of one element of a family makes the study of the other members a much easier task, since so many of the properties and chemical reactions of the elements are similar. Thus, having studied the element sulphur in some detail, it is not necessary to study selenium and tellurium so closely, for most of their properties can be predicted from the relation which they sustain to sulphur.

2. It predicts new elements. When the periodic law was first formulated there were a number of vacant places in the table which evidently belonged to elements at that time unknown. From their position in the table, Mendeléeff predicted with great precision the properties of the elements which he felt sure would one day be discovered to fill these places. Three of them, scandium, germanium, and gallium, were found within fifteen years, and their properties agreed in a remarkable way with the predictions of Mendeléeff. There are still some vacant places in the table, especially among the heavier elements.

3. It corrects errors. The physical constants of many of the elements did not at first agree with those demanded by the periodic law, and a further study of many such cases showed that errors had been made. The law has therefore done much service in indicating probable error.

Imperfections of the law. There still remain a good many features which must be regarded as imperfections in the law. Most conspicuous is the fact that the element hydrogen has no place in the table. In some of the groups elements appear in one of the families, while all of their properties show that they belong in the other. Thus sodium belongs with lithium and not with copper; fluorine belongs with chlorine and not with manganese. There are[Pg 173] two instances where the elements must be transposed in order to make them fit into their proper group. According to their atomic weights, tellurium should follow iodine, and argon should follow potassium. Their properties show in each case that this order must be reversed. The table separates some elements altogether which, in many respects have closely agreeing properties. Iron, chromium, and manganese are all in different groups, although they are similar in many respects.

The system is therefore to be regarded as but a partial and imperfect expression of some very important and fundamental relation between the substances which we know as elements, the exact nature of this relation being as yet not completely clear to us.

EXERCISES

1. Suppose that an element were discovered that filled the blank in Group O, Period 5; what properties would it probably have?

2. Suppose that an element were discovered that filled the blank in Group VI, Period 9, family B; what properties would it have?

3. Sulphur and oxygen both belong in Group VI, although in different families; in what respects are the two similar?

[Pg 174]

CHAPTER XVI

THE CHLORINE FAMILY

The family. The four elements named in the above table form a strongly marked family of elements and illustrate very clearly the way in which the members of a family in a periodic group resemble each other, as well as the character of the differences which we may expect to find between the individual members.

1. Occurrence. These elements do not occur in nature in the free state. The compounds of the last three elements of the family are found extensively 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 melting points and boiling points of the elements of the family increase with their atomic weights. A somewhat similar gradation is noted in their color and state. One atom of each of the elements combines with one atom of hydrogen to form acids, which are gases very soluble in water. The affinity of the elements for hydrogen is in[Pg 175] the inverse order of their atomic weights, fluorine having the strongest affinity and iodine the weakest. Only chlorine and iodine form oxides, and those of the former element are very unstable. The elements of the group are univalent in their compounds with hydrogen and the metals.

FLUORINE

Occurrence. The element fluorine occurs in nature most abundantly as the mineral fluorspar (CaF₂), as cryolite (Na₃AlF₆), and in the complex mineral apatite (3 Ca₃(PO₄)₂·CaF₂).

Preparation. All attempts to isolate the element resulted in failure until recent years. Methods similar to those which succeed in the preparation of the other elements of the family cannot be used; for as soon as the fluorine is liberated it combines with the materials of which the apparatus is made or with the hydrogen of the water which is always present. The preparation of fluorine was finally accomplished by the French chemist Moissan by the electrolysis of hydrofluoric acid. Perfectly dry hydrofluoric acid (HF) was condensed to a liquid and placed in a U-shaped tube made of platinum (or copper), which was furnished with electrodes and delivery tubes, as shown in Fig. 52. This liquid is not an electrolyte, but becomes such when potassium fluoride is dissolved in it. When this solution was electrolyzed hydrogen was set free at the cathode and fluorine at the anode.

Fig. 52

[Pg 176]

Properties. Fluorine is a gas of slightly yellowish color, and can be condensed to a liquid boiling at -187° under atmospheric pressure. It solidifies at -223°. It is extremely active chemically, being the most active of all the elements at ordinary temperatures.

It combines with all the common elements save oxygen, very often with incandescence and the liberation of much heat. It has a strong affinity for hydrogen and is able to withdraw it from its compounds with other elements. Because of its great activity it is extremely poisonous. Fluorine does not form any oxides, neither does it form any oxygen acids, in which respects it differs from the other members of the family.

Hydrofluoric acid (HF). Hydrofluoric acid is readily obtained from fluorspar by the action of concentrated sulphuric acid. The equation is

In its physical properties it resembles the binary acids of the other elements of this family, being, however, more easily condensed to a liquid. The anhydrous acid boils at 19° and can therefore be prepared at ordinary pressures. It is soluble in all proportions in water, and a concentrated solution—about 50%—is prepared for the market. Its fumes are exceedingly irritating to the respiratory organs, and several chemists have lost their lives by accidentally breathing them.

HENRI MOISSAN (French) (1853-1907)

Famous for his work with the electric furnace at high temperatures; prepared artificial diamonds, together with many new binary compounds such as carbides, silicides, borides, and nitrides; isolated fluorine and studied its properties and its compounds very thoroughly

[Pg 177]

Chemical properties. Hydrofluoric acid, like other strong acids, readily acts on bases and metallic oxides and forms the corresponding fluorides. It also dissolves certain metals such as silver and copper. It acts very vigorously upon organic matter, a single drop of the concentrated acid making a sore on the skin which is very painful and slow in healing. Its most characteristic property is its action upon silicon dioxide (SiO₂), with which it forms water and the gas silicon tetrafluoride (SiF₄), as shown in the equation

Glass consists of certain compounds of silicon, which are likewise acted on by the acid so that it cannot be kept in glass bottles. It is preserved in flasks made of wax or gutta-percha.

Salts of hydrofluoric acid,—fluorides. A number of the fluorides are known, but only one of them, calcium fluoride (CaF₂), is of importance. This is the well-known mineral fluorspar.

CHLORINE

Historical. While studying the action of hydrochloric acid upon the mineral pyrolusite, in 1774, Scheele obtained a yellowish, gaseous substance to which he gave a name in keeping with the phlogiston theory then current. Later it was supposed to be a compound containing oxygen. In[Pg 178] 1810, however, the English chemist Sir Humphry Davy proved it to be an element and named it chlorine.

Occurrence. Chlorine does not occur free in nature, but its compounds are widely distributed. For the most part it occurs in combination with the metals in the form of chlorides, those of sodium, potassium, and magnesium being most abundant. Nearly all salt water contains these substances, particularly sodium chloride, and very large salt beds consisting 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 warming the mineral pyrolusite (manganese dioxide, MnO₂) with concentrated hydrochloric acid. The first reaction, which seems to be similar to the action of acids upon oxides in general, is expressed in the equation

The manganese compound so formed is very unstable, however, and breaks clown according to the equation

Instead of using hydrochloric acid in the preparation of chlorine it will serve just as well to use a mixture of sodium chloride and sulphuric acid, since these two react to form hydrochloric acid. The following equations will then express the changes:

[Pg 179]

Combining these equations, the following equation expressing the complete reaction is obtained:

Since the hydrochloric acid liberated in the third equation is free to act upon manganese dioxide, it will be seen that all of the chlorine originally present in the sodium chloride is set free.

Fig. 53

2. Electrolytic method. Under the discussion of electrolysis (p. 102) it was shown that when a solution of sodium chloride is electrolyzed chlorine is evolved at the anode, while the sodium set free at the cathode reacts with the water to form hydrogen, which is evolved, and sodium hydroxide, which remains in solution. A great deal of the chlorine required in the chemical industries is now made in this way in connection with the manufacture of sodium hydroxide.

Physical properties. Chlorine is a greenish-yellow gas, which has a peculiar suffocating odor and produces a very violent effect upon the throat and lungs. Even when inhaled in small quantities it often produces all the symptoms of a[Pg 180] hard cold, and in larger quantities may have serious and even fatal action. It is quite heavy (density = 2.45) and can therefore be collected by displacement of air. One volume of water under ordinary conditions dissolves about three volumes of chlorine. The gas is readily liquefied, a pressure of six atmospheres serving to liquefy it at 0°. It forms a yellowish liquid which solidifies at -102°.

Chemical properties. At ordinary temperatures chlorine is far more active chemically than any of the elements we have so far considered, with the exception of fluorine; indeed, it is one of the most active of all elements.

1. Action on metals. A great many metals combine directly with chlorine, especially when hot. A strip of copper foil heated in a burner flame and then dropped into chlorine burns with incandescence. Sodium burns brilliantly when heated strongly in slightly moist chlorine. Gold and silver are quickly tarnished by the gas.

2. Action on non-metals. Chlorine has likewise a strong affinity for many 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 burst into flame when dropped into jars of the gas. The products formed in all cases 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 in the air continues to burn when introduced into a jar of chlorine, giving a somewhat luminous flame. A mixture of the two gases explodes violently when a spark is passed through it or when it is exposed to bright sunlight. In the latter case it is the light and not the heat which starts the action.[Pg 181]

4. Action on substances containing hydrogen. Not only will chlorine combine directly with free hydrogen but it will often abstract the element from its compounds. Thus, when chlorine is passed into a solution containing hydrosulphuric acid, sulphur is precipitated and Hydrochloric acid formed. The reaction is shown by the following equation:

With ammonia the action is similar:

The same tendency is very strikingly seen in the action of chlorine upon turpentine. The latter substance is largely made up of compounds having the composition represented by the formula C₁₀H₁₆. When a strip of paper moistened with warm turpentine is placed in a jar of chlorine dense fumes of hydrochloric acid appear and a black deposit of carbon is formed. Even water, which is a very stable compound, can be decomposed by chlorine, the oxygen being liberated. This may be shown in the following way:

Fig. 54

5. Action on color substances,—bleaching action. If strips of brightly colored cloth or some highly colored flowers are placed in quite dry chlorine, no marked change[Pg 182] in color is noticed as a rule. If, however, the cloth and flowers are first moistened, the color rapidly disappears, that is, the objects are bleached. Evidently the moisture as well as the chlorine is concerned in the action, and a study of the case shows that the chlorine has combined with the hydrogen of the water. The oxygen set free oxidizes the color substance, converting it into a colorless compound. It is evident from this explanation that chlorine will only bleach those substances which are changed into colorless compounds by oxidation.

6. Action as a disinfectant. Chlorine has also marked germicidal properties, and the free element, as well as compounds from which it is easily liberated, are used as disinfectants.

Nascent state. It will be noticed that oxygen when set free from water by chlorine is able to do what ordinary oxygen cannot do, for both the cloth and the flowers are unchanged in the air which contains oxygen. It is generally true that the activity of an element is greatest at the instant of liberation from its compounds. To express this fact elements at the instant of liberation are said to be in the nascent state. It is nascent oxygen which does the bleaching.

Hydrochloric acid (muriatic acid) (HCl). The preparation of hydrochloric acid may be discussed under two general heads:

1. Laboratory preparation. The product formed by the burning of hydrogen in chlorine is the gas hydrochloric acid. This substance is much more easily obtained, however, by treating common salt (sodium chloride) with sulphuric acid. The following equation shows the reaction:

[Pg 183]

The dry salt is placed in a flask furnished with a funnel tube and an exit tube, the sulphuric acid is added, and the flask gently warmed. The hydrochloric acid gas is rapidly given off and can be collected by displacement of air. The same apparatus can be used as was employed in the preparation of chlorine (Fig. 53).

2. Commercial preparation. Commercially, hydrochloric acid is prepared in connection with the manufacture of sodium sulphate, the reaction being the same as that just given. The reaction is carried out in a furnace, and the hydrochloric acid as it escapes in the form of gas is passed into water in which it dissolves, the solution forming the hydrochloric acid of commerce. When the materials are pure a colorless solution is obtained. The most concentrated solution has a density of 1.2 and contains 40% HCl. The commercial acid, often called muriatic acid, is usually colored yellow by impurities.

Composition of hydrochloric acid. When a solution of hydrochloric acid is electrolyzed in an apparatus similar to the one in which water was electrolyzed (Fig. 18), chlorine collects at the anode and hydrogen at the cathode. At first the chlorine dissolves in the water, but soon the water in the one tube becomes saturated with it, and if the stopcocks are left open until this is the case, and are then closed, it will be seen that the two gases are set free in equal volumes.[Pg 184]

When measured volumes of the two gases are caused to unite it is found that one volume of hydrogen combines with one of chlorine. Other experiments show that the volume of hydrochloric acid formed is just equal to the sum of the volumes of hydrogen and chlorine. Therefore one volume of hydrogen combines with one volume of chlorine to form two volumes of hydrochloric acid gas. Since chlorine is 35.18 times as heavy as hydrogen, it follows that one part of hydrogen by weight combines with 35.18 parts of chlorine to form 36.18 parts of hydrochloric acid.

Physical properties. Hydrochloric acid is a colorless gas which has an irritating effect when inhaled, and possesses a sour, biting taste, but no marked odor. It is heavier than air (density = 1.26) and is very soluble in water. Under standard conditions 1 volume of water dissolves about 500 volumes of the gas. On warming such a solution the gas escapes, until at the boiling point the solution contains about 20% by weight of HCl. Further boiling will not drive out any more acid, but the solution will distill with unchanged concentration. A more dilute solution than this will lose water on boiling until it has reached the same concentration, 20%, and will then distill unchanged. Under high pressure the gas can be liquefied, 28 atmospheres being required at 0°. Under these conditions it forms a colorless liquid which is not very active chemically. It boils at -80° and solidifies at -113°. The solution of the gas in water is used almost entirely in the place of the gas itself, since it is not only far more convenient but also more active.

Chemical properties. The most important chemical properties of hydrochloric acid are the following:

1. Action as an acid. In aqueous solution hydrochloric acid has very strong acid properties; indeed, it is one of[Pg 185] the strongest acids. It acts upon oxides and hydroxides, converting them into salts:

It acts upon many metals, forming chlorides and liberating hydrogen:

Unlike nitric and sulphuric acids it has no oxidizing action, so that when it acts on metals hydrogen is always given off.

2. Relation to combustion. Hydrochloric acid gas is not readily decomposed, and is therefore neither combustible nor a supporter of combustion.

3. Action on oxidizing agents. Although hydrochloric acid is incombustible, it can be oxidized under some circumstances, in which case the hydrogen combines with oxygen, while the chlorine is set free. Thus, when a solution of hydrochloric acid acts upon manganese dioxide part of the chlorine is set free:

Aqua regia. It has been seen that when nitric acid acts as an oxidizing agent it usually decomposes, as represented in the equation

The oxygen so set free may act on hydrochloric acid:

The complete equation therefore is

[Pg 186]

When concentrated nitric and hydrochloric acids are mixed this reaction goes on slowly, chlorine and some other substances not represented in the equation being formed. The mixture is known as aqua regia and is commonly prepared by adding one volume of nitric acid to three volumes of hydrochloric acid. It acts more powerfully upon metals and other substances than either of the acids separately, and owes its strength not to acid properties but to the action of the nascent chlorine which it liberates. Consequently, when it acts upon metals such as gold it converts them into chlorides, and the reaction can be represented by such equations as

Salts of hydrochloric acid,—chlorides. The chlorides of all the metals are known and many of them are very important compounds. Some of them are found in nature, and all can be prepared by the general method of preparing salts. Silver chloride, lead chloride, and mercurous chloride are insoluble in water and acids, and can be prepared by adding hydrochloric acid to solutions of compounds of the respective elements. While the chlorides have formulas similar to the fluorides, their properties are often quite different. This is seen in the solubility of the salts. Those metals whose chlorides are insoluble form soluble fluorides, while many of the metals which form soluble chlorides form insoluble fluorides.

Compounds of chlorine with oxygen and hydrogen. Chlorine combines with oxygen and hydrogen to form four different acids. They are all quite unstable, and most of them cannot be prepared in pure form; their salts can easily be made, however, and some of them will be met with in the[Pg 187] study of the metals. The formulas and names of these acids are as follows:

Oxides of chlorine. Two oxides are known, having the formulas Cl₂O and ClO₂. They decompose very easily and are good oxidizing agents.

BROMINE

Historical. Bromine was discovered in 1826 by the French chemist Ballard, 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 in many salt springs and salt deposits. The Stassfurt deposits in Germany and the salt waters of Ohio and Michigan are especially rich in bromides.

Preparation of bromine. The laboratory method of preparing bromine is essentially different from the commercial method.

Fig. 55

1. Laboratory method. As in the case of chlorine, bromine can be prepared by the action of hydrobromic acid (HBr) on manganese dioxide. Since hydrobromic acid is not an article of commerce, a mixture of sulphuric acid[Pg 188] and a bromide is commonly substituted for it. The materials are placed in a retort arranged as shown in Fig. 55. The end of the retort just touches the surface of the water in the test tube. On heating, the bromine distills over and is collected in the cold receiver. The equation is

2. Commercial method. Bromine is prepared commercially from the waters of salt wells which are especially rich in bromides. On passing a current of electricity through such waters the bromine is first liberated. Any chlorine liberated, however, will assist in the reaction, since free chlorine decomposes bromides, as shown in the equation

When the water containing the bromine is heated, the liberated bromine distills over into the receiver.

Physical properties. Bromine is a dark red liquid about three times as heavy as water. Its vapor has a very offensive odor and is most irritating to the eyes and throat. The liquid boils at 59° and solidifies at -7°; but even at ordinary temperatures it evaporates rapidly, forming a reddish-brown gas very similar to nitrogen peroxide in appearance. Bromine is somewhat soluble in water, 100 volumes of water under ordinary conditions dissolving 1 volume of the liquid. It is readily soluble in carbon disulphide, forming a yellow solution.

Chemical properties and uses. In chemical action bromine is very similar to chlorine. It combines directly with many of the same elements with which chlorine unites, but with less energy. It combines with hydrogen and takes away[Pg 189] the latter element from some of its compounds, but not so readily as does chlorine. Its bleaching properties are also less marked.

Bromine finds many uses in the manufacture of organic drugs and dyestuffs and in the preparation of bromides.

Hydrobromic acid (HBr). When sulphuric acid acts upon a bromide hydrobromic acid is set free:

At the same time some bromine is set free, as may be seen from the red fumes which appear, and from the odor. The explanation of this is found in the fact that hydrobromic acid is much less stable than hydrochloric acid, and is therefore more easily oxidized. Concentrated sulphuric acid is a good oxidizing agent, and oxidizes a part of the hydrobromic acid, liberating bromine:

Fig. 56

[Pg 190]

Properties. Hydrobromic acid very strikingly resembles hydrochloric acid in physical and chemical properties. It is a colorless, strongly fuming gas, heavier than hydrochloric acid and, like it, is very soluble in water. Under standard conditions 1 volume of water dissolves 610 volumes of the gas. Chemically, the chief point in which it differs from hydrochloric 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 from hydrochloric acid.

Salts of hydrobromic acid,—bromides. The bromides are very similar to the chlorides in their properties. Chlorine acts upon both bromides and free hydrobromic acid, liberating bromine from them:

Silver bromide is extensively used in photography, and the bromides of sodium and potassium are used as drugs.

IODINE

Historical. Iodine was discovered in 1812 by Courtois in the ashes of certain sea plants. Its presence was revealed by its beautiful violet vapor, and this suggested the name iodine (from the Greek for violet appearance).

Occurrence. In the combined state iodine occurs in very small quantities in sea water, from which it is absorbed by[Pg 191] certain sea plants, so that it is found in their ashes. It occurs along with bromine in salt springs and beds, and is also found in Chili saltpeter.

Preparation. Iodine may be prepared in a number of ways, the principal methods being the following:

1. Laboratory method. Iodine can readily be prepared in the laboratory from an iodide by the method used in preparing bromine, except that sodium iodide is substituted for sodium bromide. It can also be made by passing chlorine into a solution of an iodide.

Fig. 57

2. Commercial method. Commercially iodine was formerly prepared from seaweed (kelp), but is now obtained almost entirely from the deposits of Chili saltpeter. The crude saltpeter is dissolved in water and the solution evaporated until the saltpeter crystallizes. The remaining liquors, known as the "mother liquors," contain sodium iodate (NaIO₃), in which form the iodine is present in the saltpeter. The chemical reaction by which the iodine is liberated from this compound is a complicated one, depending on the fact that sulphurous acid acts upon iodic acid, setting iodine free. This reaction is shown as follows:

Physical properties. Iodine is a purplish-black, shining, heavy solid which crystallizes in brilliant plates. Even at ordinary temperatures it gives off a beautiful violet vapor, which increases in amount as heat is applied. It melts at 107° and boils at 175°. It is slightly soluble in water, but readily dissolves in alcohol, forming a brown solution (tincture of iodine), and in carbon disulphide, forming a violet solution. The element has a strong, unpleasant odor, though by no means as irritating as that of chlorine and bromine.

Chemical properties. Chemically iodine is quite similar to chlorine and bromine, but is still less active than bromine. It combines directly with many elements at ordinary temperatures. At elevated temperatures it combines with hydrogen, but the reaction is reversible and the compound formed is quite easily decomposed. Both chlorine and bromine displace it from its salts:

When even minute traces of iodine are added to thin starch paste a very intense blue color develops, and this reaction forms a delicate test for iodine. Iodine is extensively used in medicine, especially in the form of a tincture. It is also largely used in the preparation of dyes and organic drugs, iodoform, a substance used as an antiseptic, has the [Pg 193]formula CHI₃.

Hydriodic acid (HI). This acid cannot be prepared in pure condition by the action of sulphuric acid upon an iodide, since the hydriodic acid set free is oxidized by the sulphuric acid just as in the case of hydrobromic acid, but to a much greater extent. It can be prepared in exactly the same way as hydrobromic acid, iodine being substituted for bromine. It can also be prepared by passing hydrosulphuric acid into water in which iodine is suspended. The equation is

The hydriodic acid formed in this way dissolves in the water.

Properties and uses. Hydriodic acid resembles the corresponding acids of chlorine and bromine in physical properties, being a strongly fuming, colorless gas, readily soluble in water. Under standard conditions 1 volume of water dissolves about 460 volumes of the gas. It is, however, more unstable than either hydrochloric or hydrobromic acids, and on exposure to the air it gradually decomposes in accordance with the equation

Owing to the slight affinity between iodine and hydrogen the acid easily gives up its hydrogen and is therefore a strong reducing agent. This is seen in its action on sulphuric acid.

The salts of hydriodic acid, the iodides, are, in general, similar to the chlorides and bromides. Potassium iodide (KI) is the most familiar of the iodides and is largely used in medicine.

GAY-LUSSAC'S LAW OF VOLUMES

In the discussion of the composition of hydrochloric acid it was stated that one volume of hydrogen combines with one volume of chlorine to form two volumes of hydrochloric acid. With bromine and iodine similar combining ratios hold good. These facts recall the simple volume relations already noted in the study of the composition of steam and ammonia. These relations may be represented graphically in the following way:

Graph

In the early part of the past century Gay-Lussac, a distinguished French chemist, studied the volume relations of many combining gases, and concluded that similar relations always hold. His observations are summed up in the following law: When two gases combine chemically there is always a simple ratio between their volumes, and between the volume of either one of them and that of the product, provided it is a gas. By a simple ratio is meant of course the ratio of small whole numbers, as 1 : 2, 2 : 3.

EXERCISES

1. How do we account for the fact that liquid hydrofluoric acid is not an electrolyte?

2. Why does sulphuric acid liberate hydrofluoric acid from its salts?[Pg 195]

3. In the preparation of chlorine, what advantages are there in treating manganese dioxide with a mixture of sodium chloride and sulphuric acid rather 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 the bleaching 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 the strongest acids?

9. What is the meaning of the phrase aqua regia?

10. Cl₂O is the anhydride of what acid?

11. A solution of hydriodic acid on standing turns brown. How is this accounted for?

12. How can bromine vapor and nitrogen peroxide be distinguished from each other?

13. Write the equations for the reaction taking place when hydriodic acid is prepared from iodine, phosphorus, and water.

14. From their behavior toward sulphuric acid, to what class of agents do hydrobromic and hydriodic acids belong?

15. Give the derivation of the names of the elements of the chlorine family.

16. Write the names and formulas for the binary acids of the group in the 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 an iodide?

19. What weight of sodium chloride is necessary to prepare sufficient hydrochloric acid to saturate 1 l. of water under standard conditions?

20. On decomposition 100 l. of hydrochloric acid would yield how many liters of hydrogen and chlorine respectively, the gases being measured under the same conditions? Are your results in accord with the experimental facts?

[Pg 196]

CHAPTER XVII

CARBON AND SOME OF ITS SIMPLER COMPOUNDS

The family. Carbon stands at the head of a family of elements in the fourth group in the periodic table. The resemblances between the elements of this family, while quite marked, are not so striking as in the case of the elements of the chlorine family. With the exception of carbon, these elements are comparatively rare, and need not be taken up in detail in this chapter. Titanium will be referred to again in connection with silicon which it very closely resembles.

Occurrence. Carbon is found in nature in the uncombined state in several forms. The diamond is practically pure carbon, while graphite and coal are largely carbon, but contain small amounts of other substances. Its natural 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 of rocks, and are found in almost every locality. All living organisms, both plant and animal, contain a large percentage of this element, and the number of its compounds which go to make up all the vast variety of animate nature is almost limitless. Over one hundred thousand definite compounds containing carbon have been prepared. In the free state carbon occurs in three allotropic forms, two of which are crystalline and one amorphous.[Pg 197]

Crystalline carbon. Crystalline carbon occurs in two forms,—diamond and graphite.

1. Diamond. Diamonds are found in considerable quantities in several localities, especially in South Africa, the East Indies, and Brazil. The crystals belong to the regular system, but the natural stones do not show this very clearly. When found they are usually covered with a rough coating which is removed in the process of cutting. Diamond cutting is carried on most extensively in Holland.

The density of the diamond is 3.5, and, though brittle, it is one of the hardest of substances. Black diamonds, as well as broken and imperfect stones which are valueless as gems, are used for grinding hard substances. Few chemical reagents have any action on the diamond, but when heated in oxygen or the air it blackens and burns, forming carbon dioxide.

Lavoisier first showed that carbon dioxide is formed by the combustion of the diamond; and Sir Humphry Davy in 1814 showed that this is the only product of combustion, and that the diamond is pure carbon.

Fig. 58

2. Graphite. This form of carbon is found in large quantities, especially in Ceylon, Siberia, and in some localities of the United States and Canada. It is a shining black substance, very soft and greasy to the touch. Its density is about 2.15. It varies somewhat in properties according to the locality in which it is found, and is more easily attacked by reagents than is the diamond. It is also manufactured by heating carbon with a small amount of iron (3%) in an electric furnace. It is used in the manufacture of lead pencils and crucibles, as a lubricant, and as a protective covering for iron in the form of a polish or a paint.

Amorphous carbon. Although there are many varieties of amorphous carbon known, they are not true allotropic modifications. They differ merely in their degree of purity, their fineness of division, and in their mode of preparation. These substances are of the greatest importance, owing to their many uses in the arts and industries. As they occur in nature, or are made artificially, they are nearly all impure carbon, the impurity depending on the particular substance in question.

1. Pure carbon. Pure amorphous carbon is best prepared by charring sugar. This is a substance consisting of carbon, hydrogen, and oxygen, the latter two elements being present in the ratio of one oxygen atom to two of hydrogen.[Pg 199] When sugar is strongly heated the oxygen and hydrogen are 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 from vast accumulations of vegetable matter in former ages, which became covered over with earthy material and were thus protected from rapid decay. Under various natural agencies the organic matter was slowly changed into coal. In anthracite these changes have gone the farthest, and this variety of coal is nearly pure carbon. Soft or bituminous coals contain considerable organic matter besides carbon and mineral substances. When heated strongly out of contact with air the organic matter is decomposed and the resulting volatile matter is driven off in the form of gases and vapors, and only the mineral matter and carbon remain behind. The gaseous product is chiefly illuminating gas and the solid residue is coke. Some of the coke is found as a dense cake on the sides and roof of the retort. This is called retort carbon and is quite pure.

3. Charcoal. This is prepared from wood in the same way that coke is made from coal. When the process is carried on in retorts the products expelled by the heat are saved. Among these are many valuable substances such as wood alcohol and acetic acid. Where timber is abundant the process is carried out in a wasteful way, by merely covering piles of wood with sod and setting the wood on fire. Some wood burns and the heat from this decomposes the wood not burned, forming charcoal from it. The charcoal, of course, contains the mineral part of the wood from which it is formed.[Pg 200]

4. Bone black. This is sometimes called animal charcoal, and is made by charring bones and animal refuse. The organic part of the materials is thus decomposed and carbon is left in a very finely divided state, scattered through the mineral part which consists largely of calcium phosphate. For some uses this mineral part is removed by treatment with hydrochloric acid and prolonged washing.

5. Lampblack. Lampblack and soot are products of imperfect combustion of 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 contains various oily materials.

Properties. While the various forms of carbon differ in many properties, especially in color and hardness, yet they are all odorless, tasteless solids, insoluble in water and characterized by their stability towards heat. Only in the intense heat of the electric arc does carbon volatilize, passing directly from the solid state into a vapor. Owing to this fact the inside surface of an incandescent light bulb after being used for some time becomes coated with a dark film of carbon. It is not acted on at ordinary temperatures by most reagents, but at a higher temperature it combines directly with many of the elements, forming compounds called carbides. When heated in the presence of sufficient oxygen it burns, forming carbon dioxide.

Uses of carbon. The chief use of amorphous carbon is for fuel to furnish heat and power for all the uses of civilization. An enormous quantity of carbon in the form of the purer coals, coke, and charcoal is used as a reducing agent in the manufacture of the various metals, especially in the metallurgy of iron. Most of the metals are found in nature as oxides, or in forms which can readily be[Pg 201] converted into oxides. When these oxides are heated with carbon the oxygen is abstracted, leaving the metal. Retort carbon and coke are used to make electric light carbons and battery plates, while lampblack is used for indelible inks, printer's ink, and black varnishes. Bone black and charcoal have the property of absorbing large volumes of certain gases, as well as smaller amounts of organic matter; hence they are used in filters to remove noxious gases and objectionable colors and odors from water. Bone black is used extensively in the sugar refineries to remove coloring matter from the impure sugars.

Chemistry of carbon compounds. Carbon is remarkable for the very large number of compounds which it forms with the other elements, especially with oxygen and hydrogen. Compounds containing carbon are more numerous than all others put together, and the chemistry of these substances presents peculiarities not met with in the study of other substances. For these reasons the systematic study of carbon compounds, or of organic chemistryas it is usually called, must be deferred until the student has gained some knowledge of the chemistry of other elements. An acquaintance with a few of the most familiar carbon compounds is, however, essential for the understanding of the general principles of chemistry.

Compounds of carbon with hydrogen,—the hydrocarbons. Carbon unites with hydrogen to form a very large number of compounds called hydrocarbons. Petroleum and natural gas are essentially mixtures of a great variety of these hydrocarbons. Many others are found in living plants, and still others are produced by the decay of organic matter in the absence of air. Only two of them, methane and acetylene, will be discussed here.[Pg 202]

Methane (marsh gas) (CH₄). This is one of the most important of these hydrocarbons, and constitutes about nine tenths of natural gas. As its name suggests, it is formed in marshes by the decay of vegetable matter under water, and bubbles of the gas are often seen to rise when the dead leaves on the bottom of pools are stirred. It also collects in mines, and, when mixed with air, is called fire damp by the miners because 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 in closed vessels, and is therefore a principal constituent of coal gas.

Preparation. Methane is prepared in the laboratory by heating sodium or calcium acetate with soda-lime. Equal weights of fused sodium acetate and soda-lime are thoroughly dried, then mixed and placed in a good-sized, hard-glass test tube fitted with a one-holed stopper and delivery tube. The mixture is gradually heated, and when the air has been displaced from the tube the gas is collected in bottles by displacement of water. Soda-lime is a mixture of sodium and calcium hydroxides. Regarding it as sodium hydroxide alone, the equation is

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 use of miners, to prevent the dreadful mine explosions then common, due to methane mixed with air. The invention consisted in surrounding the upper part of the common miner's lamp with a mantle of wire gauze and the lower part with glass (Fig. 59). It has been seen that two gases will not combine until raised to their[Pg 203] kindling temperature, and if while combining they are cooled below this point, the combination ceases. A flame will not pass through a wire gauze because the metal, being a good conductor of heat, takes away so much heat from the flame that the gases are cooled below the kindling temperature. When a lamp so protected is brought into an explosive mixture the gases inside the wire mantle burn in a series of little explosions, giving warning to the miner that the air is unsafe.

Fig. 59

Acetylene (C₂H₂). This is a colorless gas usually having a disagreeable odor due to impurities. It is now made in large quantities from calcium carbide (CaC₂). This substance is formed when coal and lime are heated together in an electric furnace. When treated with water the carbide is decomposed, yielding acetylene:

Under ordinary conditions the gas burns with a very smoky flame; in burners constructed so as to secure a large amount of oxygen it burns with a very brilliant white light, and hence is used as an illuminant.

Laboratory preparation. The gas can be prepared readily in a generator such as is shown in Fig. 60. The inner tube contains fragments of calcium carbide, while the outer one is filled with water. As long as the stopcock is closed the water cannot rise in the inner tube. When the stopcock is open the water rises, and, coming into contact with the carbide in the inner tube, generates acetylene. This escapes through the stopcock, and after the air has been expelled may be lighted as it issues from the burner.

Fig. 60

Carbon forms two oxides, namely, carbon dioxide (CO₂) and carbon monoxide (CO).[Pg 204]

Carbon dioxide (CO₂). Carbon dioxide is present in the air to the extent of about 3 parts in 10,000, and this apparently small amount is of fundamental importance in nature. In some localities it escapes from the earth in great quantities, and many spring waters carry large amounts of it in solution. When these highly charged spring waters reach the surface of the earth, and the pressure on them is removed, the carbon dioxide escapes with effervescence. It is a product of the oxidation of all organic matter, and is therefore formed in fires as well as in the process of decay. It is thrown off from the lungs of all animals in respiration, and is a product of many fermentation processes such as vinegar making and brewing. Combined with metallic oxides it forms vast deposits of carbonates in nature.

Preparation. In the laboratory carbon dioxide is always prepared by the action of an acid upon a carbonate, usually calcium carbonate, the apparatus shown in Fig. 39 serving the purpose very well. This reaction might be expected to produce carbonic acid, thus:

Carbonic acid is very unstable, however, and decomposes into its anhydride, CO₂, and water, thus:

The complete reaction is represented by the equation

Physical properties. Carbon dioxide is a colorless, practically odorless gas whose density is 1.5. Its weight may be inferred from the fact that it can be siphoned, or poured like water, from one vessel downward into another. At 15°[Pg 205] and under ordinary pressure it dissolves in its own volume of water and imparts a somewhat biting, pungent taste to it. It is easily condensed, and is now prepared commercially in this form by pumping the gas into steel cylinders (see Fig. 6) which are kept cold during the process. When the liquid is permitted to escape into the air part of it instantly evaporates, and in so doing absorbs so much heat that another portion is solidified, the solid form strikingly resembling snow in appearance. This snow is very cold and mercury can easily be frozen with it.

Solid carbon dioxide. Cylinders of liquid carbon dioxide are inexpensive, and should be available in every school. To demonstrate the properties of solid carbon dioxide, the cylinder should be placed across the table and supported in such a way that the stopcock end is several inches lower than the other end. A loose bag is made by holding the corners of a handkerchief around the neck of the stopcock, and the cock is then turned on so that the gas rushes out in large quantities. Very quickly a considerable quantity of the snow collects in the handkerchief. To freeze mercury, press a piece of filter paper into a small evaporating dish and pour the mercury upon it. Coil a flat spiral upon the end of a wire, and dip the spiral into the mercury. Place a quantity 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 may be removed from the dish by the wire serving as a handle. The filter paper is to prevent the mercury from sticking to the dish; the ether dissolves the solid carbon dioxide and promotes its rapid conversion into gas.

Chemical properties. Carbon dioxide is incombustible, since it is, like water, a product of combustion. It does not support combustion, as does nitrogen peroxide, because the oxygen in it is held in very firm chemical union with the carbon. Very strong reducing agents, such as highly heated carbon, can take away half of its oxygen:

[Pg 206]

Uses. The relation of carbon dioxide to plant life has been discussed in a previous chapter. Water highly charged with carbon dioxide is used for making soda water and similar beverages. Since it is a non-supporter of combustion and can be generated readily, carbon dioxide is also used as a fire extinguisher. Some of the portable fire extinguishers are simply devices for generating large amounts of the gas. It is not necessary that all the oxygen should be kept away from the fire in order to smother it. A burning candle is extinguished in air which contains only 2.5% of carbon dioxide.

Carbonic acid (H₂CO₃). Like most of the oxides of the non-metallic elements, carbon dioxide is an acid anhydride. It combines with water to form an acid of the formula H₂CO₃, called carbonic acid:

The acid is, however, very unstable and cannot be isolated. Only a very small amount of it is actually formed when carbon dioxide is passed into water, 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, and these are quite stable:

Action of carbon dioxide on bases. This conduct is explained by the principles of reversible reactions. The equation

is a reversible equation, and the extent to which the reaction progresses depends upon the relative concentrations of each of the three factors in it. Equilibrium is ordinarily reached when very little H₂CO₃ is formed. If a base is present in the water to combine with the H₂CO₃ as fast as it is formed, all of the CO₂ is converted [Pg 207]into H₂CO₃, and thence into a carbonate.

Salts of carbonic acid,—carbonates. The carbonates form a very important 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 action of carbon dioxide on solutions of the bases, as has just been explained.

The insoluble carbonates are formed as precipitates when soluble salts are treated with a solution of a soluble carbonate. Thus the insoluble calcium carbonate can be made by bringing together solutions of calcium chloride and sodium carbonate:

Most of the carbonates are decomposed by heat, yielding an oxide of the metal and carbon dioxide. Thus lime (calcium oxide) is made by strongly heating calcium carbonate:

Acid carbonates. Like all acids containing two acid hydrogen atoms, carbonic acid can form both normal and acid salts. The acid carbonates are made by treating a normal carbonate with an excess of carbonic acid. With few exceptions they are very unstable, heat decomposing them even when in solution.

Action of carbon dioxide on calcium hydroxide. If carbon dioxide is passed into clear lime water, calcium carbonate is at first precipitated:

Advantage is taken of this reaction in testing for the presence of carbon dioxide, as already explained in the chapter on the atmosphere. If the current of carbon dioxide is continued, the precipitate[Pg 208] soon dissolves, because the excess of carbonic acid forms calcium acid carbonate which is soluble:

If now the solution is heated, the acid carbonate is decomposed and calcium carbonate once more precipitated:

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 is conducted over highly heated carbon, the monoxide is formed, thus:

It is therefore often formed in stoves when the air draught is insufficient. Water gas, which contains large amounts of carbon monoxide, is made by partially oxidizing carbon with steam:

2. By the partial reduction of carbon dioxide. When carbon dioxide is conducted over highly heated carbon it is reduced to carbon monoxide by the excess of carbon:

When coal is burning in a stove or grate carbon dioxide is at first formed in the free supply of air, but as the hot gas rises through the glowing coal it is reduced to carbon monoxide. When the carbon monoxide reaches the free air above the coal it takes up oxygen to form carbon dioxide, 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 carbon monoxide is usually prepared by the action of[Pg 209] concentrated sulphuric acid upon oxalic acid. The latter substance has the formula C₂H₂O₄. The sulphuric acid, owing to its affinity for water, decomposes the oxalic acid, as represented in the equation

Properties. Carbon monoxide is a light, colorless, almost odorless gas, very difficult to liquefy. Chemically it is very active, combining directly with a great many substances. It has a great affinity for oxygen and is therefore combustible and a good reducing agent. Thus, if carbon monoxide is passed over hot copper oxide, the copper is reduced to the metallic state:

When inhaled it combines with the red coloring matter of the blood and in this way prevents the absorption of oxygen, so that even a small quantity of the gas may prove fatal.

Fig. 61

[Pg 210]

Carbon disulphide (CS₂). Just as carbon combines with oxygen to form carbon dioxide, so it combines with sulphur to form carbon disulphide (CS₂). This compound has been described in the chapter on sulphur.

Hydrocyanic acid (prussic acid)(HCN). Under the proper conditions carbon unites with nitrogen and hydrogen to form the acid HCN, called hydrocyanic acid. It is a weak, volatile acid, and is therefore easily prepared by treating its salts with sulphuric acid:

It is most familiar as a gas, though it condenses to a colorless liquid boiling at 26°. It has a peculiar odor, suggesting bitter almonds, and is extremely poisonous either when inhaled or when taken into the stomach. 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 of sodium and potassium being the best known. These are white solids and are extremely poisonous.

Solutions of potassium cyanide are alkaline. A solution of potassium cyanide 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 is nevertheless measurably dissociated. Hydrocyanic acid is one of the weakest of acids, and dissociates[Pg 211] to an extremely slight extent. When a cyanide such as potassium cyanide dissolves it freely dissociates, and the CN^- ions must come to an equilibrium with the H⁺ ions derived from the water:

The result of this equilibrium is that quite a number of H⁺ ions from the water are converted into undissociated HCN molecules. But for every H⁺ ion so removed an OH^- ion remains free, and this will give the solution alkaline properties.

EXERCISES

1. How can you prove that the composition of the different allotropic forms of carbon is the same?

2. Are lampblack and bone black allotropic forms of carbon? Will equal amounts of heat be liberated in the combustion of 1 g. of each?

3. How could you judge of the relative purity of different forms of carbon?

4. Apart from its color, why should carbon be useful in the preparation of 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 of carbon dioxide in water?

8. What compound would be formed by passing carbon dioxide into a solution of ammonium hydroxide? Write the equation.

9. Write equations for the preparation of K₂CO₃; of BaCO₃; of MgCO₃.

10. In what respects are carbonic and sulphurous acids similar?

11. Give three reasons why the reaction which takes place when a solution 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?[Pg 212]

14. Could a solution of sodium hydroxide be substituted for the solution of calcium hydroxide in testing for carbon dioxide?

15. What weight of sodium hydroxide is necessary to neutralize the carbonic acid formed by the action of hydrochloric acid on 100 g. of calcium carbonate?

16. What weight of calcium carbonate would be necessary to prepare sufficient carbon dioxide to saturate 10 l. of water at 15° and under ordinary 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. of acetylene measured at 20° and 740 mm.?

18. How would the volume of a definite amount of carbon monoxide compare with the volume of carbon dioxide formed by its combustion, the measurements being made under the same conditions?

[Pg 213]

CHAPTER XVIII

FLAMES,—ILLUMINANTS

Conditions necessary for flames. It has been seen that when two substances unite chemically, with the production of light and heat, the act of union is called combustion. When one of the substances undergoing combustion remains solid at the temperature occasioned by the combustion, light may be given off, but there is no flame. Thus iron wire burning in oxygen throws off a shower of sparks and is brilliantly incandescent, but no flame is seen. When, however, both of the substances are gases or vapors at the temperature reached in the combustion, the act of union is accompanied by a flame.

Flames from burning liquids or solids. Many substances which are liquids or solids at ordinary temperatures burn with a flame because the heat of combustion vaporizes them slowly, and the flame is due to the union of this vapor with the gas supporting the combustion.

Supporter of combustion. That gas which surrounds the flame and constitutes the atmosphere in which the combustion occurs is said to support the combustion. The other gas which issues into this atmosphere is said to be the combustible gas. Thus, in the ordinary combustion of coal gas in the air the coal gas is said to be combustible, while the air is regarded as the supporter of combustion. These terms are entirely relative, however, for a jet of air issuing into an atmosphere of coal gas will burn when ignited, the coal gas supporting the combustion.[Pg 214] Ordinarily, when we say that a gas is combustible we mean that it is combustible in an atmosphere of air.

Fig. 62

Either gas may be the supporter of combustion. That the terms combustible and supporter of combustion are merely relative may be shown in the following way: A lamp chimney A is fitted with a cork and glass tubes, as shown in Fig. 62. The tube C should have a diameter of from 12 to 15 mm. A thin sheet of asbestos in which is cut a circular opening 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 gas is admitted to the chimney through the tube B. The air in the chimney is soon expelled through the tube C, and the gas itself is then lighted at the lower end of this tube. The hand is now removed from the opening in the asbestos, when the flame at the end of the tube at once rises 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 and may be lighted. The flame at the top of the asbestos board is due to the combustion of coal gas in air, while the flame within the chimney is due to the combustion of air in coal gas, the air being drawn up through the tube by the escaping gas.

Appearance of flames. The flame caused by the union of hydrogen and oxygen is almost colorless and invisible. Chlorine and hydrogen combine with a pale violet flame, carbon monoxide burns in oxygen with a blue flame, while ammonia burns with a deep yellow flame. The color and appearance of flames are therefore often quite characteristic of the particular combustion which occasions them.

Structure of flames. When the gas undergoing combustion issues from a round opening into an atmosphere of the gas supporting combustion, as is the case with the burning Bunsen burner (Fig. 63), the flame is generally[Pg 215] conical in outline. It consists of several distinct cones, one within the other, the boundary between them being marked by differences of color or luminosity. In the simplest flame, of which hydrogen burning in oxygen is a good example, these cones are two in number,—an inner one, formed by unburned gas, and an outer one, usually more or less luminous, consisting of the combining gases. This outer one is in turn surrounded by a third envelope of the products of combustion; this envelope is sometimes invisible, as in the present case, but is sometimes faintly luminous. The lower part of the inner cone of the flame is quite cool and consists of unburned gas. Toward the top of the inner cone the gas has become heated to a high temperature by the burning envelope surrounding it. On reaching the supporter of combustion on the outside it is far above its kindling temperature, and combustion follows with the evolution of much heat. The region of combustion just outside the inner cone is therefore the hottest part of the flame.

Fig. 63

Oxidizing and reducing flames. Since the tip of the outside cone consists of very hot products of combustion mixed with oxygen from the air, a substance capable of oxidation placed in this part of the flame becomes very hot and is easily oxidized. The oxygen with which it combines comes, of course, from the atmosphere, and not from the products 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. This region consists of a highly heated combustible gas, which has not yet reached a supply of oxygen.[Pg 216]

If a substance rich in oxygen, such as a metallic oxide, is placed in this region of the flame, the heated gases combine with its oxygen and the substance is reduced. This part of the flame is called the reducing flame. These flames are used in testing certain substances, especially minerals. For this purpose they are produced by blowing into a small luminous Bunsen flame from one side through a blowpipe. This is a tube of the shape shown in Fig. 64. The flame is directed in any desired way and has the oxidizing and reducing regions very clearly marked (Fig. 65). It is non-luminous from the same causes which render the open Bunsen burner flame non-luminous, the gases from the lungs serving to furnish oxygen and to dilute the combustible gas.

Fig. 64

Fig. 65

Luminosity of flames. The luminosity of flames is due to a number of distinct causes, and may therefore be increased or diminished in several ways.

1. Presence of solid matter. The most obvious of these causes is the presence in the flame of incandescent solid matter. Thus chalk dust sifted into a non-luminous flame renders it luminous. When hydrocarbons form a part of the combustible gas, as they do in nearly all illuminating gases and oils, some carbon is usually set free in the process of combustion. This is made very hot by the flame and becomes incandescent, giving out light. In a well-regulated flame it is afterward burned up, but when the supply of oxygen is insufficient it escapes from the flame as lampblack or soot. That it is temporarily present in a well-burning luminous flame may be demonstrated by holding a cold object, such as a small evaporating dish, in the flame for a few seconds. This cold object cools the carbon below its kindling temperature, and it is deposited on the object as soot.[Pg 217]

2. Pressure. A second factor in the luminosity of flames is the pressure under which the gases are burning. Under increased pressure there is more matter in a given volume of a gas, and the chemical action is more energetic than when the gases are rarefied. Consequently there is more heat and light. A candle burning on a high mountain gives less light than when it burns at the sea level.

If the gas is diluted with a non-combustible gas, the effect is the same as if it is rarefied, for under these conditions there is less combustible gas in a given volume.

3. Temperature. The luminosity also depends upon the temperature attained in the combustion. In general the hotter the flame the greater the luminosity; hence cooling the gases before combustion diminishes the luminosity of the flame they will make, because it diminishes the temperature attained in the combustion. Thus the luminosity of the Bunsen 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 and cooled by the air drawn in. The oxygen thus introduced into the flame also causes the combustion of the hot particles of carbon which would otherwise 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 use for the production of light and heat. The three chief varieties are coal gas, water gas, and natural gas. The use of acetylene gas has already been referred to.

Coal gas. Coal gas is made by heating bituminous coal in large retorts out of contact with the air. Soft or bituminous coal contains, in addition to large amounts of carbon, considerable quantities of compounds of hydrogen, oxygen, nitrogen, and sulphur. When distilled the nitrogen is liberated partly in the form of ammonia and cyanides and partly as free nitrogen gas; the sulphur is converted into hydrogen sulphide, carbon disulphide, and oxides of sulphur; the oxygen into water and oxides of carbon. The[Pg 218] remaining hydrogen is set free partly as hydrogen and partly in combination with carbon in the form of hydrocarbons. The most important of these is methane, with smaller quantities of many others, some of which are liquids or solids at ordinary temperatures. The great bulk of the carbon remains behind as coke and retort carbon.

The manufacture of coal gas. In the manufacture of coal gas it is necessary to separate from the volatile constituents formed by the heating of the coal all those substances which are either solid or liquid 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 in Fig. 66. The coal is heated in air-tight retorts illustrated by A. The volatile products escape through the pipe X and bubble into the tarry liquid in the large pipe B, known as the hydraulic main, which runs at right angles to the retorts. Here is deposited the greater portion of the solid and liquid products, forming a tarry mass known as coal tar. Much of the ammonia also remains dissolved in this liquid. The partially purified gas then passes into the pipes C, which serve to cool it and further remove the solid and liquid matter. The gas then passes into D, which is filled with coke over which a jet of water is sprayed. The water still further cools the gas and at the same time partially removes such gaseous products as hydrogen sulphide and ammonia, which are soluble 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 much of the carbon dioxide present. From E the gas passes into the large gas holder F, from which it is distributed through pipes to the places where it is burned.

Fig. 66

[Pg 219]

Water gas. Water gas is essentially a mixture of carbon monoxide and hydrogen. It is made by passing steam over very hot anthracite coal, when the reaction shown in the following equation takes place:

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 burn with non-luminous flames. This is accomplished by passing it into heaters containing highly heated petroleum oils. The gas takes up hydrocarbon 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 and hydrogen burn with very hot flames. It has little odor and is very poisonous. Its use is therefore attended with some risk, since leaks in pipes are very likely to escape notice.

Natural gas. This substance, so abundant in many localities, varies much in composition, but is composed principally of methane. When used for lighting purposes it is usually burned in a burner resembling an open Bunsen, the illumination being furnished by an incandescent mantle. This is the case in the familiar Welsbach burner. Contrary to statements frequently made, natural gas contains no free hydrogen.[Pg 220]

TABLE SHOWING COMPOSITION OF GASES

Fuels. A variety of substances are used as fuels, the most important of them being wood, coal, and the various gases mentioned above. Wood consists mainly of compounds of carbon, hydrogen, and oxygen. The composition of coal and the fuel gases has been given. Since these fuels are composed principally of carbon and hydrogen or their compounds, the chief products of combustion are carbon dioxide and water. The practice of heating rooms with portable gas or oil stoves with no provision for removing the products of combustion is to be condemned, since the carbon dioxide is generated in sufficient quantities to render the air unfit for breathing. Rooms so heated also become very damp from the large amount of water vapor formed in the combustion, and which in[Pg 221] cold weather condenses on the window glass, causing the glass to "sweat." Both coal and wood contain a certain amount of mineral substances which constitute the ashes.

The electric furnace. In recent years electric furnaces have come into wide use in operations requiring a very high temperature. Temperatures as high as 3500° can be easily reached, whereas the hottest oxyhydrogen flame is not much above 2000°. These furnaces are constructed on one of two general principles.

Fig. 67

1. Arc furnaces. In the one type the source of heat is an electric arc formed between carbon electrodes separated a little from each other, as shown in Fig. 67. The substance to be heated is placed in a vessel, usually a graphite crucible, just below the arc. The electrodes and crucible are surrounded by materials which fuse with great difficulty, such as magnesium oxide, the walls of the furnace being so shaped as to reflect the heat downwards upon the contents of the crucible.

Fig. 68

2. Resistance furnaces. In the other type of furnace the heat is generated by the resistance offered to the current in its passage through 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, are introduced through the sides of the furnace. The materials to be heated, C, are filled into the furnace up to the electrodes, and a layer of broken 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 through the broken coke the electrical current encounters great resistance. This generates great heat, and the charge surrounding the coke is brought to a very high temperature. The advantage of this type of furnace is that the temperature can be regulated to any desired intensity.[Pg 222]

EXERCISES

1. Why does charcoal usually burn with no flame? How do you account for the 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 lamp possess?

4. (a) In what respects does the use of the Welsbach mantle resemble that of lime in the calcium light? (b) If the mantle were made of carbon, 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 the combustion 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 sulphur compounds 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 a more intense heat?

10. What volume of oxygen is necessary to burn 100 l. of marsh gas and what volume of carbon dioxide would be formed, all of the gases being measured under standard conditions?

11. Suppose a cubic meter of Pennsylvania natural gas, measured under standard conditions, were to be burned. How much water by weight would result?

[Pg 223]

CHAPTER XIX

MOLECULAR WEIGHTS, ATOMIC WEIGHTS, FORMULAS

Introduction. In the chapter on The Atomic Theory, it was shown that if it were true that two elements uniting to form a compound always combined in the ratio of one atom of one element to one atom of the other element, it would be a very easy matter to decide upon figures which would represent the relative weights of the different atoms. It would only be necessary to select some one element as a standard and determine the weight of every element which combines with a definite weight (say 1 g.) of the standard element. The figures so obtained would evidently represent the relative weights of the atoms.

But the law of multiple proportion at once reminds us that two elements may unite in several proportions; and there is no simple way to determine the number of atoms present in the molecule of any compound. Consequently the problem of deciding upon the relative atomic weights is not an easy one. To the solution of this problem we must now turn.

Dalton's method of determining atomic weights. When Dalton first advanced the atomic theory he attempted to solve this problem by very simple methods. He thought that when only one compound of two elements is known it is reasonable to suppose that it contains one atom of each element. He therefore gave the formula HO to water, and HN to ammonia. When more than two compounds were known he assumed that the most familiar or the most stable one had the simple formula. He then determined the atomic weight as[Pg 224] explained above. The results he obtained were contradictory and very far from satisfactory, and it was soon seen that some other method, resting on much more scientific grounds, must be found to decide what compounds, if any, have a single atom of each element present.

Determination of atomic weights. Three distinct steps are involved in the determination of the atomic weight of an element: (1) determination of the equivalent, (2) determination of molecular weights of its compounds, and (3) deduction of the exact atomic weight from the equivalent and molecular weights.

1. Determination of the equivalent. By the equivalent of an element is meant the weight of the element which will combine with a fixed weight of some other element chosen as a standard. It has already been explained that oxygen has been selected as the standard element for atomic weights, with a weight of 16. This same standard will serve very well as a standard for equivalents. The equivalent of an element is the weight of the element which will combine with 16 g. of oxygen. Thus 16 g. of oxygen combines with 16.03 g. of sulphur, 65.4 g. of zinc, 215.86 g. of silver, 70.9 g. of chlorine. These figures, therefore, represent the equivalent weights of these elements.

Relation of atomic weights to equivalents. According to the atomic theory combination always takes place between whole numbers of atoms. Thus one atom unites with one other, or with two or three; or two atoms may unite with three, or three with five, and so on.

When oxygen combines with zinc the combination must be between definite numbers of the two kinds of atoms. Experiment shows that these two elements combine in the ratio of 16 g. of oxygen to 65.4 g. of zinc. If one atom of[Pg 225] oxygen combines with one atom of zinc, then this ratio must be the ratio between the weights of the two atoms. If one atom of oxygen combines with two atoms of zinc, then the ratio between the weights of the two atoms will be 16: 32.7. If two atoms of oxygen combine with one atom 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 must be some multiple or submultiple of the equivalent; in other words, the equivalent multiplied by 1/2, 1, 2, or 3 will give the atomic weight.

Combining weights. A very interesting relation holds good between the equivalents of the various elements. We have just seen that the figures 16.03, 65.4, 215.86, and 70.9 are the equivalents respectively of sulphur, zinc, silver, and chlorine. These same figures represent the ratios by weight in which these elements combine among themselves. Thus 215.86 g. of silver combine with 70.9 g. of chlorine and with 2 × 16.03 g. 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 obtained for each element which will represent its combining value, and for this reason is called its combining weight. It is important to notice that the fact that a combining weight can be obtained for each element is not a part of a theory, but is the direct result of experiment.

Elements with more than one equivalent. It will be remembered that oxygen combines with hydrogen in two ratios. In one case 16 g. of oxygen combine with 2.016 g. of hydrogen to form water; in the other 16 g. of oxygen combine with 1.008 g. of hydrogen to form hydrogen dioxide. The equivalents of hydrogen are therefore 2.016 and 1.008. Barium combines with 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 16 g. of oxygen to 68.7 g. of barium.[Pg 226]

In each case one equivalent is a simple multiple of the other, so the fact 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 of the equivalent.

2. The determination of molecular weights. To decide the question as to which multiple of the equivalent correctly represents the atomic weight of an element, it has been found necessary to devise a method of determining the molecular weights of compounds containing the element in question. Since the molecular weight of a compound is merely the sum of the weights of all the atoms present in it, it would seem to be impossible to determine the molecular weight of a compound without first knowing the atomic weights of the constituent atoms, and how many atoms of each element are present in the molecule. But certain facts have been discovered which suggest a way in which this can be done.

Avogadro's hypothesis. We have seen that the laws of Boyle, Charles, and Gay-Lussac apply to all gases irrespective of their chemical character. This would lead to the inference that the structure of gases must be quite simple, and that it is much the same in all gases.

In 1811 Avogadro, an Italian physicist, suggested that if we assume all gases under the same conditions of temperature and pressure to have the same number of molecules in a given volume, we shall have a probable explanation of the simplicity of the gas laws. It is difficult to prove the truth of this hypothesis by a simple experiment, but there are so many facts known which are in complete harmony with this suggestion that there is little doubt that it expresses the truth. Avogadro's hypothesis may be stated thus: Equal volumes of all gases under the same conditions of temperature and pressure contain the same number of molecules.

[Pg 227]

Avogadro's hypothesis and molecular weights. Assuming that Avogadro's hypothesis is correct, we have a very simple means for deciding upon the relative weights of molecules; for if equal volumes of two gases contain the same number of molecules, the weights of the two volumes must be in the same ratio as the weights of the individual molecules which they contain. If we adopt some one gas as a standard, we can express the weights of all other gases as compared with this one, and the same figures will express the relative weights of the molecules of which the gases are composed.

Oxygen as the standard. It is important that the same standard should be adopted for the determination of molecular weights as has been decided upon for atomic weights and equivalents, so that the three values may be in harmony with each other. Accordingly it is best to adopt oxygen as the standard element with which to compare the molecular weights of other gases, being careful to keep the oxygen atom equal to 16.

The oxygen molecule contains two atoms. One point must not be overlooked, however. We desire to have our unit, the oxygen atom, equal to 16. The method of comparing the weights of gases just suggested compares the molecules of the gases with the molecule of oxygen. Is the molecule and the atom of oxygen the same thing? This question is answered by the following considerations.

We have seen that when steam is formed by the union of oxygen and hydrogen, two volumes of hydrogen combine with one volume of oxygen to form two volumes of steam. Let us suppose that the one volume of oxygen contains 100[Pg 228] molecules; then the two volumes of steam must, according to Avogadro's hypothesis, contain 200 molecules. But each of these 200 molecules must contain at least one atom of oxygen, or 200 in all, and these 200 atoms came from 100 molecules of oxygen. It follows that each molecule of oxygen must contain at least two atoms of oxygen.

Evidently this reasoning merely shows that there are at least two atoms in the oxygen molecule. There may be more than that, but as there is no evidence to this effect, we assume that the molecule contains two atoms only.

It is evident that if we wish to retain the value 16 for the atom of oxygen we must take twice this value, or 32, for the value of the oxygen molecule, when using it as a standard for molecular weights.

Determination of the molecular weights of gases from their weights compared with oxygen. Assuming the molecular weight of oxygen to be 32, Avogadro's hypothesis gives us a ready means for determining the molecular weight of any other gas, for all that is required is to know its 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 as 1 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 in grams corresponding to the molecular weight of the gas, the calculation is much simplified. It has been found that 32 g. of oxygen, under standard conditions, measure 22.4 l. This same volume of hydrogen weighs 2.019 g.; of chlorine 70.9 g.; of hydrochloric acid 36.458 g. The weights of these equal volumes must be proportional to their molecular weights, and since[Pg 229] the weight of the oxygen is the same as the value of its molecular weight, so too will the weights of the 22.4 l. of the other gases be equal to the value of their molecular weights.

As a summary we can then make the following statement: The molecular weight 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 actual experiment it is easier to determine the density of a gas than the weight of a definite volume of it. The density of a gas is usually defined as its weight compared with that of an equal volume of air. Having determined the density of a gas, its weight compared with oxygen may be determined by multiplying its density by the ratio between the weights of air and oxygen. This ratio is 0.9046. To compare it with our standard for atomic weights we must further multiply it by 32, since the standard is 1/32 the weight of oxygen molecules. The steps then are these:

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 atomic weights.

We have, then, the formula:

or, still more briefly,

The value found by this method for the determination of molecular weights will of course agree with those found[Pg 230] by calculating the weight of 22.4 l. of the gas, since both methods depend on the same principles.

Fig. 69

3. Deduction of atomic weights from molecular weights and equivalents. We have now seen how the equivalent of an element and the molecular weight of compounds containing the element can be obtained. Let us see how it is possible to decide which multiple of the equivalent really is the true atomic weight. As an example, let us suppose that the equivalent of nitrogen has been found to be 7.02 and that it is desired to obtain its atomic weight. The next step is to obtain the molecular weights of a large number of compounds containing nitrogen. The following will serve:

[Pg 231]

Method of calculation. The densities of the various gases in the first column of this table are determined by experiment, and are fairly accurate but not entirely so. By multiplying these densities by 28.9 the molecular weights of the compounds as given in the second column are obtained. By chemical analysis it is possible to determine the percentage composition of these substances, and the percentages of nitrogen in them as determined by analysis are given in the third column. If each of these molecular weights is multiplied in turn by the percentage of nitrogen in the compound, the product will be the weight of the nitrogen in the molecular weight of the compound. This will be the sum of the weights of the nitrogen atoms in the molecule. These values are given in the fourth column in the table.

If a large number of compounds containing nitrogen are studied in this way, it is probable that there will be included in the list at least one substance whose molecule contains a single nitrogen atom. In this case the number in the fourth column will be the approximate atomic weight of nitrogen. On comparing the values for nitrogen in the table it will be seen that a number which is approximately 14 is the smallest, and that the 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 be determined very accurately, and consequently the part in them due to nitrogen is a little uncertain, as will be seen in the table. All we can tell by this method is that the true weight is very near 14. The equivalent can however be determined very accurately, and we have seen that it is some multiple or submultiple[Pg 232] of the true atomic weight. Since molecular-weight determinations have shown that in the case of nitrogen the atomic weight is near 14, and we have found the equivalent to be 7.02, it is evident that the true atomic weight is twice the equivalent, or 7.02 × 2 = 14.04.

Summary. These, then, are the steps necessary to establish the atomic weight of an element.

1. Determine the equivalent accurately by analysis.

2. Determine the molecular weight of a large number of compounds of the element, and by analysis the part of the molecular weight due to the element. The smallest number so obtained will be approximately the atomic weight.

3. Multiply the equivalent by the small whole number (usually 1, 2, or 3), which will make a number very close to the approximate atomic weight. The figure so obtained will be the true atomic weight.

Molecular weights of the elements. It will be noticed that the molecular weight 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 gaseous compounds is 14.04. The simplest explanation that can be given for this is that the gaseous nitrogen is made up of molecules, each of which contains two atoms. In this respect it resembles oxygen; for we have seen that an entirely different line of reasoning leads us to believe that the molecule of oxygen contains two atoms. When we wish to indicate molecules of these gases the symbols N₂ and O₂ should be used. When we desire to merely show the weights taking part in a reaction this is not necessary.

The vapor densities of many of the elements show that, like oxygen and nitrogen, their molecules consist of two atoms. In other cases, particularly among the metals,[Pg 233] 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 has led to the conclusion that it has three. The formation of ozone from oxygen can therefore be represented by the equation

Other methods of determining molecular weights. It will be noticed that Avogadro's law gives us a method by which we can determine the relative weights of the molecules of two gases because it enables us to tell when we are dealing with an equal number of the two kinds of molecules. If by any other means we can get this information, we can make use of the knowledge so gained to determine the molecular weights of the two substances.

Raoult's laws. Two laws have been discovered which give us just such information. They are known as Raoult's laws, and can be stated as follows:

1. When weights of substances which are proportional to their molecular weights are dissolved in the same weight of solvent, the rise of the boiling point is the same in each case.

2. When weights of substances which are proportional to their molecular weights are dissolved in the same weight of solvent, the lowering of the freezing point is the same in each case.

By taking advantage of these laws it is possible to determine when two solutions contain the same number of molecules of two dissolved substances, and consequently the relative molecular weights of the two substances.

Law of Dulong and Petit. In 1819 Dulong and Petit discovered a very interesting relation between the atomic[Pg 234] weight of an element and its specific heat, which holds true for elements in the solid state. If equal weights of two solids, say, lead and silver, are heated through the same range of temperature, as from 10° to 20°, it is found that very different amounts of heat are required. The amount of heat required to change the temperature of a solid or a liquid by a definite amount compared with the amount required to change the temperature of an equal weight of water by the same amount is called its specific heat. Dulong and Petit discovered the following law: The specific heat of an element in the solid form multiplied by its atomic weight is approximately equal to the constant 6.25. That is,

Consequently,

This law is not very accurate, but it is often possible by means of it to decide upon what multiple of the equivalent is the real atomic weight. Thus the specific heat of iron is found by experiment to be 0.112, 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 with molecular weights to consider how the formula of a compound is decided upon, for the two subjects are very closely associated. Some examples will make clear the method followed.

The molecular weight of a substance containing hydrogen and chlorine was 36.4. By analysis 36.4 parts of the substance was found to contain 1 part of hydrogen and 35.4 parts of chlorine. As these are the simple atomic[Pg 235] weights of the two elements, the formula of the compound must be HCl.

A substance consisting of oxygen and hydrogen was found to have a molecular weight of 34. Analysis showed that in 34 parts of the substance there were 2 parts of hydrogen and 32 parts of oxygen. Dividing these figures by the atomic weights of the two elements, we get 2 ÷ 1 = 2 for H; 32 ÷ 16 = 2 for O. The formula is therefore H₂O₂.

A substance containing 2.04% H, 32.6% S, and 65.3% O was found to have a molecular 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% = 64 parts of O. If the molecule weighs 98, the hydrogen atoms present must together weigh 2, the sulphur atoms 32, and the oxygen atoms 64. Dividing these figures by the respective atomic weights of the three elements, we have, for H, 2 ÷ 1 = 2 atoms; for S, 32 ÷ 32 = 1 atom; for O, 64 ÷ 16 = 4 atoms. Hence the formula is H₂SO₄.

We have, then, this general procedure: Find the percentage composition of the substance and also its molecular weight. Multiply the molecular weight successively by the percentage of each element present, to find the amount of the element in the molecular weight of the compound. The figures so obtained will be the respective parts of the molecular weight due to the several atoms. Divide by the atomic weights of the respective elements, and the quotient will be the number of atoms present.

Avogadro's hypothesis and chemical calculations. This law simplifies many chemical calculations.

1. Application to volume relations in gaseous reactions. Since equal volumes of gases contain an equal number of[Pg 236] molecules, it follows that when an equal number of gaseous molecules of two or more gases take part in a reaction, the reaction will involve equal volumes of the gases. In the equation

since 1 molecule of each of the gases CO₂ and CO is set free from each molecule of oxalic acid, the two substances must always be set free in equal volumes.

Acetylene burns in accordance with the equation

Hence 2 volumes of acetylene will react with 5 volumes of oxygen to form 4 volumes of carbon dioxide and 2 volumes of steam. That the volume relations may be correct a gaseous element must be given its molecular formula. Thus oxygen must be written O₂ and not 2O.

2. Application to weights of gases. It will be recalled that the molecular weight of a gas is determined by ascertaining the weight of 22.4 l. of the gas. This weight in grams is called the gram-molecular weight of a gas. If the molecular weight of any gas is known, the weight of a liter of the gas under standard conditions may be determined by dividing its gram-molecular weight by 22.4. Thus the gram-molecular weight of a hydrochloric acid gas is 36.458. A liter of the gas will therefore weigh 36.458 ÷ 22.4 = 1.627 g.

EXERCISES

1. From the following data calculate the atomic weight of sulphur. The equivalent, as obtained by an analysis of sulphur dioxide, is 16.03. The densities and compositions of a number of compounds containing sulphur are as follows:[Pg 237]

2. Calculate the formulas for compounds of the following compositions:

3. The molecular weight of ammonia is 17.06; of sulphur dioxide is 64.06; of chlorine is 70.9. From the molecular weight calculate the weight of 1 l. of each of these gases. Compare your results with the table on the back cover of the book.

4. From the molecular weight of the same gases calculate the density of each, referred to air as a standard.

5. A mixture of 50 cc. of carbon monoxide and 50 cc. of oxygen was exploded in a eudiometer, (a) What gases remained in the tube after the explosion? (b) What was the volume of each?

6. In what proportion must acetylene and oxygen be mixed to produce the greatest 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 these specific heats calculate the atomic weights of each of the elements.

[Pg 238]

CHAPTER XX

THE PHOSPHORUS FAMILY

The family. The elements constituting this family belong in the same group with nitrogen and therefore resemble it in a general way. They exhibit a regular gradation of physical properties, as is shown in the above table. The same general gradation is also found in their chemical properties, phosphorus being an acid-forming element, while bismuth is essentially a metal. The other two elements are intermediate in properties.

Compounds. In general the elements of the family form compounds having similar composition, as is shown in the following table:

In the case of phosphorus, arsenic, and antimony the oxides are acid anhydrides. Salts of at least four acids of each of these three elements are known, the free acid in[Pg 239] some instances being unstable. The relation of these acids to the corresponding anhydrides may be illustrated as follows, phosphorus being taken as an example:

PHOSPHORUS

History. The element phosphorus was discovered by the alchemist Brand, of Hamburg, in 1669, while searching for the philosopher's stone. Owing to its peculiar properties and the secrecy which was maintained about its preparation, it remained a very rare and costly substance until the demand for it in the manufacture of matches brought about its production on a large scale.

Occurrence. Owing to its great chemical activity phosphorus never occurs free in nature. In the form of phosphates it is very abundant and widely distributed. Phosphorite and sombrerite are mineral forms of calcium phosphate, while apatite consists of calcium phosphate together with calcium fluoride or chloride. These minerals form very large deposits and are extensively mined for use as fertilizers. Calcium phosphate is a constituent of all fertile soil, having been supplied to the soil by the disintegration of rocks containing it. It is the chief mineral constituent of bones of animals, and bone ash is therefore nearly pure calcium phosphate.

Preparation. Phosphorus is now manufactured from bone ash or a pure mineral phosphate by heating the phosphate with sand and carbon in an electric furnace. The materials[Pg 240] are fed in at M (Fig. 70) by the feed screw F. The phosphorus vapor escapes at P and is condensed under water, while the calcium silicate is tapped off as a liquid at S. The phosphorus obtained in this way is quite impure, and is purified by distillation.

Fig. 70

Physical properties. The purified phosphorus is a pale yellowish, translucent, waxy solid which melts at 43.3° and boils at 269°. It can therefore be cast into any convenient form under warm water, and is usually sold in the market in the form of sticks. It is quite soft and can be easily cut with a knife, but this must always be done while the element is covered with water, since it is extremely inflammable, and the friction of the knife blade is almost[Pg 241] sure to set it on fire if cut in the air. It is not soluble in water, but is freely soluble in some other liquids, notably in carbon disulphide. Its density is 1.8.

Chemical properties. Exposed to the air phosphorus slowly combines with oxygen, and in so doing emits a pale light, or phosphorescence, which can be seen only in a dark place. The heat of the room may easily raise the temperature to the kindling point of phosphorus, when it burns with a sputtering flame, giving off dense fumes of oxide of phosphorus. It burns with dazzling brilliancy in oxygen, and combines directly with many other elements, especially with sulphur and the halogens. On account of its great affinity for oxygen it is always preserved under water.

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 used as a poison for rats and other vermin.

Red phosphorus. On standing, yellow phosphorus gradually undergoes a remarkable change, being converted into a dark red powder which has a density of 2.1. It no longer takes fire easily, neither does it dissolve in carbon disulphide. It is not poisonous and, in fact, seems to be an entirely different substance. The velocity of this change increases with rise in temperature, and the red phosphorus is therefore prepared by heating the yellow just below the boiling point (250°-300°). When distilled and quickly condensed the red form changes back to the yellow. This is in accordance with the general rule that when a substance capable[Pg 242] of existing in several allotropic forms is condensed from a gas or crystallized from the liquid state, the more unstable variety forms first, and this then passes into the more stable forms.

Compounds of phosphorus with hydrogen. Phosphorus forms several compounds with hydrogen, the best known of which is phosphine (PH₃) analogous to ammonia (NH₃).

Preparation of phosphine. Phosphine is usually made by heating phosphorus with a strong solution of potassium hydroxide, the reaction being a complicated one.

Fig. 71

Properties. Phosphine is a gas of unpleasant odor and is exceedingly poisonous. Like ammonia it forms salts with the halogen acids. Thus we have phosphonium chloride (PH₄Cl) analogous to ammonium chloride (NH₄Cl). The phosphonium salts are of but little importance.

Oxides of phosphorus. Phosphorus forms two well-known oxides,—the trioxide (P₂O₃) and the pentoxide (P₂O₅), sometimes called phosphoric anhydride. When phosphorus burns in an insufficient supply of air the product is partially the trioxide; in oxygen or an excess of air the pentoxide is formed. The pentoxide is much the better known of the two. It is a snow-white, voluminous powder whose most marked property is its great attraction for water. It has no chemical action upon most gases, so that they can be very thoroughly dried by allowing them to pass through properly arranged vessels containing phosphorus pentoxide.

Acids of phosphorus. The important acids of phosphorus are the following:

These may be regarded as combinations of the oxides of phosphorus with water according to the equations given in the discussion of the characteristics of the family.[Pg 244]

1. Phosphorous acid (H₃PO₃). Neither the acid nor its salts are at all frequently met with in chemical operations. It can be easily obtained, however, in the form of transparent crystals when phosphorus trichloride is treated with water and the resulting solution is evaporated:

Its most interesting property is its tendency to take up oxygen and pass over into phosphoric acid.

2. Orthophosphoric acid (phosphoric acid) (H₃PO₄). This acid can be obtained by dissolving phosphorus pentoxide in boiling water, as represented in the equation

It is usually made by treating calcium phosphate with concentrated sulphuric acid. The calcium sulphate produced in the reaction is nearly insoluble, and can be filtered off, leaving the phosphoric acid in solution. Very pure acid is made by oxidizing phosphorus with nitric acid. It forms large colorless crystals which are exceedingly soluble in water. Being a tribasic acid, it forms acid as well as normal salts. Thus the following compounds of sodium are known:

These salts are sometimes called respectively primary, secondary, and tertiary phosphates. They may be prepared by bringing together phosphoric acid and appropriate quantities of sodium hydroxide. Phosphoric acid also forms mixed salts, that is, salts containing two different metals. The most familiar compound of this kind is microcosmic [Pg 245]salt, which has the formula Na(NH₄)HPO₄.

Orthophosphates. The orthophosphates form an important class of salts. The normal salts are nearly all insoluble and many of them occur in nature. The secondary phosphates are as a rule insoluble, while most of the primary salts are soluble.

3. Pyrophosphoric acid (H₄P₂O₇). On heating orthophosphoric acid to about 225° pyrophosphoric acid is formed in accordance with the following equation:

It is a white crystalline solid. Its salts can be prepared by heating a secondary phosphate:

4. Metaphosphoric acid (glacial phosphoric acid) (HPO₃). This acid is formed when orthophosphoric acid is heated above 400°:

It is also formed when phosphorus pentoxide is treated with cold water:

It is a white crystalline solid, and is so stable towards heat that it can be fused and even volatilized without decomposition. On cooling from the fused state it forms a glassy solid, and on this account is often called glacial phosphoric acid. It possesses the property of dissolving small quantities of metallic oxides, with the formation of compounds which, in the case of certain metals, have characteristic colors. It is therefore used in the detection of these metals.

While the secondary phosphates, on heating, give salts of pyrophosphoric acid, the primary phosphates yield salts of metaphosphoric acid. The equations representing these reactions are as follows:

Fertilizers. When crops are produced year after year on the same field certain constituents of the soil essential to plant growth are removed, and the soil becomes impoverished and unproductive. To make the land once more[Pg 246] fertile these constituents must be replaced. The calcium phosphate of the mineral deposits or of bone ash serves well as a material for restoring phosphorus to soils exhausted of that essential element; but a more soluble substance, which the plants can more readily assimilate, is desirable. It is better, therefore, to convert the insoluble calcium phosphate into the soluble primary phosphate before it is applied as fertilizer. It will be seen by reference to the formulas for the orthophosphates (see page 244) that in a primary phosphate only one hydrogen atom of phosphoric acid is replaced by a metal. Since the calcium atom always replaces two hydrogen atoms, it might be thought that there could be no primary calcium phosphate; but if the calcium atom replaces one hydrogen atom from each of two molecules of phosphoric acid, the salt Ca(H₂PO₄)₂ will result, and this is a primary phosphate. It can be made by treatment of the normal phosphate with the necessary amount of sulphuric acid, calcium sulphate being formed at the same time, thus:

The resulting mixture is a powder, which is sold as a fertilizer under the name of "superphosphate of lime."

ARSENIC

Occurrence. Arsenic occurs in considerable quantities in nature as the native element, as the sulphides realgar (As₂S₂) and orpiment (As₂S₃), as oxide (As₂O₃), and as a constituent of many metallic sulphides, such as arsenopyrite (FeAsS).

Preparation. The element is prepared by purifying the native arsenic, or by heating the arsenopyrite in iron tubes,[Pg 247] out of contact with air, when the reaction expressed by the following equation occurs:

The arsenic, being volatile, condenses in chambers connected with the heated tubes. It is also made from the oxide by reduction with carbon:

Properties. Arsenic is a steel-gray, metallic-looking substance of density 5.73. Though resembling metals in appearance, it is quite brittle, being easily powdered in a mortar. When strongly heated it sublimes, that is, it passes into a vapor without melting, and condenses again to a crystalline solid when the vapor is cooled. Like phosphorus it can be obtained in several allotropic forms. It alloys readily with some of the metals, and finds its chief use as an alloy with lead, which is used for making shot, the alloy being harder than pure lead. When heated on charcoal with the blowpipe it is converted into an oxide which volatilizes, leaving the charcoal unstained by any oxide coating. It burns readily in chlorine gas, forming arsenic trichloride,—

Unlike most of its compounds, the element itself is not poisonous.

Arsine (AsH₃). When any compound containing arsenic is brought into the presence of nascent hydrogen, arsine (AsH₃), corresponding to phosphine and ammonia, is formed. The reaction when oxide of arsenic is so treated is

[Pg 248]

Arsine is a gas with a peculiar garlic-like odor, and is intensely poisonous. A single bubble of pure gas has been known to prove fatal. It is an unstable compound, decomposing into its elements when heated to a moderate temperature. It is combustible, burning with a pale bluish-white flame to form arsenic trioxide and water when air is in excess:

When the supply of air is deficient water and metallic arsenic are formed:

These reactions make the detection of even minute quantities of arsenic a very easy problem.

Fig. 72

Oxides of arsenic. Arsenic forms two oxides, As₂O₃ and As₂O₅, corresponding to those of phosphorus. Of these arsenious oxide, or arsenic trioxide (As₂O₃), is much better known, and is the substance usually called white arsenic, or merely arsenic. It is found as a mineral, but is usually obtained as a by-product in burning pyrite in the sulphuric-acid industry. The pyrite has a small amount of arsenopyrite in it, and when this is burned arsenious oxide is formed as a vapor together with sulphur dioxide:

The arsenious oxide is condensed in appropriate chambers. It is a rather heavy substance, obtained either as a crystalline powder or as large, vitreous lumps, resembling lumps of porcelain in appearance. It is very poisonous, from 0.2 to 0.3 g. being a fatal dose. It is frequently given as 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, and hence is absorbed slowly by the system. Arsenious oxide is also used as a chemical reagent in glass making and in the dye industry.[Pg 250]

Acids of arsenic. Like the corresponding oxides of phosphorus, the oxides of arsenic are acid anhydrides. In solution they combine with bases to form salts, corresponding to the salts of the acids of phosphorus. Thus we have salts of the following acids:

Several other acids of arsenic are also known. Not all of these can be obtained as free acids, since they tend to lose water and form the oxides. Thus, instead of obtaining arsenious acid (H₃AsO₃), the oxide As₂O₃ is obtained:

Salts of all the acids are known, however, and some of them have commercial value. Most of them are insoluble, and some of the copper salts, which are green, are used as pigments. Paris green, which has a complicated formula, is a well-known insecticide.

Antidote for arsenical poisoning. The most efficient antidote for arsenic poisoning is ferric hydroxide. It is prepared as needed, according to the equation

Sulphides of arsenic. When hydrogen sulphide is passed into an acidified solution containing an arsenic compound the arsenic is precipitated as a bright yellow sulphide, thus:

In this respect arsenic resembles the metallic elements, many of which produce sulphides under similar conditions. The sulphides of arsenic, both those produced artificially and those found in nature, are used as yellow pigments.

ANTIMONY

Occurrence. Antimony occurs in nature chiefly as the sulphide (Sb₂S₃), called stibnite, though it is also found as oxide and as a constituent of many complex minerals.[Pg 251]

Preparation. Antimony is prepared from the sulphide in a very simple manner. The sulphide is melted with scrap iron in a furnace, when the iron combines with the sulphur to form a slag, or liquid layer of melted iron sulphide, while the heavier liquid, antimony, settles to the bottom and is drawn off from time to time. The reaction involved is represented by the equation

Physical properties. Antimony is a bluish-white, metallic-looking substance whose density is 6.7. It is highly crystalline, hard, and very brittle. It has a rather low melting point (432°) and expands very noticeably on solidifying.

Chemical properties. In chemical properties antimony resembles arsenic in many particulars. It forms the oxides Sb₂O₃ and Sb₂O₅, and in addition Sb₂O₄. It combines with the halogen elements with great energy, burning brilliantly in chlorine to form antimony trichloride (SbCl₃). When heated on charcoal with the blowpipe it is oxidized and forms a coating of antimony oxide on the charcoal which has a characteristic bluish-white color.

Stibine (SbH₃). The gas stibine (SbH₃) is formed under conditions which are very similar to those which produce arsine, and it closely resembles the latter compound, though it is still less stable. It is very poisonous.

Metallic properties of antimony. The physical properties of the element are those of a metal, and the fact that its sulphide is precipitated by hydrogen sulphide shows that it acts like a metal in a chemical way. Many other reactions show that antimony has more of the properties of a metal than of a non-metal. The compound Sb(OH)₃, corresponding to arsenious acid, while able to act as a weak acid is also able to act as a weak base with strong acids. For example, when treated with concentrated hydrochloric acid antimony chloride is formed:

A number of elements act in this same way, their hydroxides under some conditions being weak acids and under others weak bases.

ALLOYS

Some metals when melted together thoroughly intermix, and on cooling form a homogeneous, metallic-appearing substance called an alloy. Not all metals will mix in this way, and in some cases definite chemical compounds are formed and separate out as the mixture solidifies, thus destroying the uniform quality of the alloy. In general the melting point of the alloy is below the average of the melting points of its constituents, and it is often lower than any one of them.

Antimony forms alloys with many of the metals, and its chief commercial use is for such purposes. It imparts to its alloys high density, rather low melting point, and the[Pg 253] property of expanding on solidification. Such an alloy is especially useful in type founding, where fine lines are to be reproduced on a cast. Type metal consists of antimony, lead, and tin. Babbitt metal, used for journal bearings in machinery, contains the same metals in a different proportion together with a small percentage 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 commerce comes from Saxony, and from Mexico and Colorado, but it is not an abundant element.

Preparation. It is prepared by merely heating the ore containing the native bismuth and allowing the melted metal to run out into suitable vessels. Other ores are converted into oxides and reduced by heating with carbon.

Physical properties. Bismuth is a heavy, crystalline, brittle metal nearly the color of silver, but with a slightly rosy tint which distinguishes 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 ordinary temperatures.

Chemical properties. When heated with the blowpipe on charcoal, bismuth gives a coating of the oxide Bi₂O₃. This has a yellowish-brown color which easily distinguishes it from the oxides formed by other metals. It combines very readily with the halogen elements, powdered bismuth burning readily in chlorine. It is not very easily acted upon by hydrochloric acid, but nitric and sulphuric acids act upon it in the same 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[Pg 254] of these melt in hot water. For example, Wood's metal, consisting of bismuth, lead, tin, and cadmium, melts at 60.5°.

Compounds of bismuth. Unlike the other elements of this group, bismuth has almost no acid properties. Its chief oxide, Bi₂O₃, is basic in its properties. It dissolves in strong acids and forms salts of bismuth:

The nitrate and chloride of bismuth can be obtained as well-formed colorless crystals. When treated with water the salts are decomposed in the manner explained in the following paragraph.

HYDROLYSIS

Many salts such as those of antimony and bismuth form solutions which are somewhat acid in reaction, and must therefore contain hydrogen ions. This is accounted for by the same principle suggested to explain the fact 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 a very slight extent. When Bi⁺⁺⁺ ions from bismuth chloride, which dissociates very readily, are brought in contact with the OH^- ions from water, the two come to the equilibrium expressed in the equation

For every hydroxyl ion removed from the solution in this way a hydrogen ion is left free, and the solution becomes acid in reaction.[Pg 255]

Reactions of this kind and that described under potassium cyanide are called hydrolysis.

DEFINITION: Hydrolysis is the action of water upon a salt to form an acid and a base, one of which is very slightly dissociated.

Conditions favoring hydrolysis. While hydrolysis is primarily due to the slight extent to which either the acid or the base formed is dissociated, several other factors have an influence upon the extent to which it will take place.

1. Influence of mass. Since hydrolysis is a reversible reaction, the relative masses of the reacting substances influence the point at which equilibrium will be reached. In the equilibrium

the addition of more water will result in the formation of more bismuth hydroxide and hydrochloric acid. The addition of more hydrochloric acid will convert some of the bismuth hydroxide into bismuth chloride.

2. Formation of insoluble substances. When one of the products of hydrolysis is nearly insoluble in water the solution will become saturated with it as soon as a very little has been formed. All in excess of this will precipitate, and the reaction will go on until the acid set free increases sufficiently to bring about an equilibrium. Thus a considerable amount of bismuth and antimony hydroxides are precipitated when water is added to the chlorides of these elements. The greater the dilution the more hydroxide precipitates. The addition of hydrochloric acid in considerable quantity will, however, redissolve the precipitate.

Partial hydrolysis. In many cases the hydrolysis of a salt is only partial, resulting in the formation of basic salts instead of the free base. Most of these basic salts are insoluble in water, which accounts for their ready formation. Thus bismuth chloride may hydrolyze by successive steps, as shown in the equations

The basic salt so formed may also lose water, as shown in the equation

[Pg 256]

The salt represented in the last equation is sometimes called bismuth oxychloride, or bismuthyl chloride. The corresponding nitrate, BiONO₃, is largely used in medicine under the name of subnitrate of bismuth. In these two compounds the group of atoms, BiO, acts as a univalent metallic radical and is called bismuthyl. Similar basic salts 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 and iodine? Write the equations showing the action of water on these compounds.

3. In the preparation of phosphine, why is coal gas passed into the flask? What other gases would serve the same purpose?

4. Give the formula for the salt which phosphine forms with hydriodic acid. Give the name of the compound.

5. Could phosphoric acid be substituted for sulphuric acid in the preparation of the common acids?

6. Write the equations for the preparation of the three sodium salts of orthophosphoric acid.

7. Why does a solution of disodium hydrogen phosphate react alkaline?

8. On the supposition that bone ash is pure calcium phosphate, what weight of it would be required in the preparation of 1 kg. of phosphorus?

9. If arsenopyrite is heated in a current of air, what products are formed?

10. (a) Write equations for the complete combustion of hydrosulphuric acid, methane, and arsine. (b) In what respects are the reactions similar?

11. Write the equations for all the reactions involved in Marsh's test for 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 the phosphorus family?

[Pg 257]

CHAPTER XXI

SILICON, TITANIUM, BORON

General. Each of the three elements, silicon, titanium, and boron, belongs to a separate periodic family, but they occur near together in the periodic grouping and are very similar in both physical and chemical properties. Since the other elements in their families are either so rare that they cannot be studied in detail, or are best understood in connection with other elements, it is convenient to consider these three together at this point.

The three elements are very difficult to obtain in the free state, owing to their strong attraction for other elements. They can be prepared by the action of aluminium or magnesium on their oxides and in impure state by reduction with carbon in an electric furnace. They are very hard and melt only at the highest temperatures. At ordinary temperatures they are not attacked by oxygen, but when strongly heated they burn with great brilliancy. Silicon and boron are not attacked by acids under ordinary conditions; titanium is easily dissolved by them.[Pg 258]

SILICON

Occurrence. Next to oxygen silicon is the most abundant element. It does not occur free in nature, but its compounds are very abundant and of the greatest importance. It occurs almost entirely in combination with oxygen as silicon dioxide (SiO₂), often called silica, or with oxygen and various metals in the form of salts of silicic acids, or silicates. These compounds form a large fraction of the earth's crust. Most plants absorb small amounts of silica from the soil, and it is also found in minute quantities in animal organisms.

Preparation. The element is most easily prepared by reducing pure powdered quartz with magnesium powder:

Properties. As would be expected from its place in the periodic table, silicon resembles carbon in many respects. It can be obtained in several allotropic forms, corresponding to those of carbon. The crystallized form is very hard, and is inactive toward reagents. The amorphous variety has, in general, properties more similar to charcoal.

Compounds of silicon with hydrogen and the halogens. Silicon hydride (SiH₄) corresponds in formula to methane (CH₄), but its properties are more like those of phosphine (PH₃). It is a very inflammable gas of disagreeable odor, and, as ordinarily prepared, takes fire spontaneously on account of the presence of impurities.

Silicon combines with the elements of the chlorine family to form such compounds as SiCl₄ and SiF₄. Of these silicon fluoride is the most familiar and interesting. As stated in the discussion of fluorine, it is formed when[Pg 259] hydrofluoric acid acts upon silicon dioxide or a silicate. With silica the reaction is thus expressed:

It is a very volatile, invisible, poisonous gas. In contact with water it is partially decomposed, as shown in the equation

The hydrofluoric acid so formed combines with an additional amount of silicon fluoride, forming the complex fluosilicic acid (H₂SiF₆), thus:

Silicides. As the name indicates, silicides are binary compounds consisting of silicon and some other element. They are very stable at high temperatures, and are usually made by heating the appropriate substances in an electric furnace. The most important one is carborundum, which is a silicide of carbon of the formula CSi. It is made by heating coke and sand, which is a form of silicon dioxide, in an electric furnace, the process being extensively carried on at Niagara Falls. The following equation represents the reaction

The substance so prepared consists of beautiful purplish-black crystals, which are very hard. Carborundum is used as an abrasive, that is, as a material for grinding and polishing very hard substances. Ferrosilicon is a silicide of iron alloyed with an excess of iron, which finds extensive use in the manufacture of certain kinds of steel.[Pg 260]

Manufacture of carborundum. The mixture of materials is heated in a large resistance furnace for about thirty-six hours. After the reaction is completed there is left a core of graphite G. Surrounding this core is a layer of crystallized carborundum C, about 16 in. thick. Outside this is a shell of amorphous carborundum A. The remaining materials M are unchanged and are used for a new charge.

Fig. 73

Silicon dioxide (silica) (SiO₂). This substance is found in a great variety of forms in nature, both in the amorphous and in the crystalline condition. In the form of quartz it is found in beautifully formed six-sided prisms, sometimes of great size. When pure it is perfectly transparent and colorless. Some colored varieties are given special names, as amethyst (violet), rose quartz (pale pink), smoky or milky quartz (colored and opaque). Other varieties of silicon dioxide, some of which also contain water, are chalcedony, onyx, jasper, opal, agate, and flint. Sand and sandstone are largely silicon dioxide.

Properties. As obtained by chemical processes silicon dioxide is an amorphous white powder. In the crystallized state it is very hard and has a density of 2.6. It is insoluble in water and in most chemical reagents, and requires the hottest oxyhydrogen flame for fusion. Acids, excepting hydrofluoric acid, have little action on it, and it requires the most energetic reducing agents to deprive it of oxygen. It is the anhydride of an acid, and consequently it dissolves in fused alkalis to form silicates. Being nonvolatile, it will drive out most other anhydrides when heated[Pg 261] to a high temperature with their salts, especially when the silicates so formed are fusible. The following equations illustrate this property:

Silicic acids. Silicon forms two simple acids, orthosilicic acid (H₄SiO₄) and metasilicic acid (H₂SiO₃). Orthosilicic acid is formed as a jelly-like mass when orthosilicates are treated with strong acids such as hydrochloric. On attempting to dry this acid it loses water, passing into metasilicic or common silicic acid:

Metasilicic acid when heated breaks up into silica and water, thus:

Polysilicic acids. Silicon has the power to form a great many complex acids which may be regarded as derived from the union of several molecules of the orthosilicic acid, with the loss of water. Thus we have

These acids cannot be prepared in the pure state, but their salts form many of the crystalline rocks in nature. Feldspar, for example, has the formula KAlSi₃O₈, and is a mixed salt of the acid H₄Si₃O₈, whose formation is represented in the equation above. Kaolin has the formula Al₂Si₂O₇·2H₂O. Many other examples will be met in the study of the metals.[Pg 262]

Glass. When sodium and calcium silicates, together with silicon dioxide, are heated to a very high temperature, the mixture slowly fuses to a transparent liquid, which on cooling passes into the solid called glass. Instead of starting with sodium and calcium silicates it is more convenient and economical to heat sodium carbonate (or sulphate) and lime with an excess of clean sand, the silicates being formed during the heating:

Fig. 74

The mixture is heated below the fusing point for some time, so that the escaping carbon dioxide may not spatter the hot liquid; the heat is then increased and the mixture kept in a state of fusion until all gases formed in the reaction have escaped.

Molding and blowing of glass. The way in which the melted mixture is handled in the glass factory depends upon the character of the article to be made. Many articles, such as bottles, are made by blowing the plastic glass into hollow molds of the desired shape. The mold is first opened, as shown in Fig. 74. A lump of plastic glass A on the hollow rod B is lowered into the mold, which is then closed by the handles C. By blowing into the tube the glass is blown into the shape of the mold. The mold is then opened and the bottle lifted out. The neck of the bottle must be cut off at the proper place and the sharp edges rounded off in a flame.

Other objects, such as lamp chimneys, are made by getting a lump of plastic glass on the end of a hollow iron rod and blowing it into the desired shape without the help of a mold, great skill being required in the manipulation of the glass. Window glass is made by blowing large hollow cylinders about 6 ft. long and 1-1/2 ft. in diameter. These are cut longitudinally, and are then placed in an oven and heated until they soften, when they are flattened out into plates (Fig. 75). Plate glass is cast into flat slabs, which are then ground and polished to perfectly plane surfaces.[Pg 263]

Varieties of glass. The ingredients mentioned above make a soft, easily fusible glass. If potassium carbonate is substituted for the sodium carbonate, the glass is much harder and less easily fused; increasing the amount of sand has somewhat the same effect. Potassium glass is largely used in making chemical glassware, since it resists the action of reagents better than the softer sodium glass. If lead oxide is substituted 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 optical instruments and artificial jewels.

Fig. 75

Coloring of glass. Various substances fused along with the glass mixture give characteristic colors. The amber color of common bottles is due to iron compounds in the glass; in other cases iron colors the glass green. Cobalt compounds color it deep blue; those of manganese give it an amethyst tint and uranium compounds impart a peculiar yellowish green color. Since iron is nearly always present in the ingredients, glass is usually slightly yellow. This color can be removed by adding the proper amount of manganese dioxide, for the amethyst color of manganese and the yellow of iron together produce white light.

Nature of glass. Glass is not a definite chemical compound and its composition varies between wide limits. Fused glass is really a solution of various silicates, such as those of calcium and lead, in fused sodium or potassium silicate. A certain amount of silicon dioxide is also present. This solution is then allowed to solidify under such conditions of cooling that the dissolved substances do not separate from the solvent. The compounds which are used to color the glass are sometimes converted into silicates, which then dissolve in the glass, giving it a uniform color. In other cases, as in the milky glasses which resemble porcelain in appearance, the color or opaqueness is due to the finely divided color material evenly distributed throughout the glass, but not dissolved in it. Milky glass is made by mixing calcium fluoride, tin oxide, or some other insoluble substance in the melted glass. Copper or gold in metallic form scattered through glass gives it shades of red.[Pg 264]

TITANIUM

BORON

Occurrence. Boron is never found free in nature. It occurs as boric acid (H₃BO₃), and in salts of polyboric acids, which usually have very complicated formulas.

Preparation and properties. Boron can be prepared from its oxide by reduction with magnesium, exactly as in the case of silicon. It resembles silicon very strikingly in its properties. It occurs in several allotropic forms, is very hard when crystallized, and is rather inactive toward reagents. It forms a hydride, BH₃, and combines directly with the elements of the chlorine family. Boron fluoride (BF₃) is very similar to silicon fluoride in its mode of formation and chemical properties.

Boric oxide (B₂O₃). Boron forms one well-known oxide, B₂O₃, called boric anhydride. It is formed as a glassy mass by heating boric acid to a high temperature. It absorbs water very readily, uniting with it to form boric acid again:

In this respect it differs from silicon dioxide, which will not combine directly with water.[Pg 265]

Boric acid (H₃BO₃). This is found in nature in considerable quantities and forms one of the chief sources of boron compounds. It is found dissolved in the water of hot springs in some localities, particularly in Italy. Being volatile with steam, the vapor which escapes from these springs has some boric acid in it. It is easily obtained from these sources by condensation and evaporation, the necessary 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, hot solution of borax with sulphuric acid. Boric acid being sparingly soluble in water crystallizes out on cooling:

The substance is a mild antiseptic, and on this account is often used in medicine and as a preservative for canned foods and milk.

Metaboric and polyboric acids. When boric acid is gently heated it is converted into metaboric acid (HBO₂):

On heating metaboric acid to a somewhat higher temperature tetraboric acid (H₂B₄O₇) is formed:

Many other complex acids of boron are known.

Borax. Borax is the sodium salt of tetraboric acid, having the formula Na₂B₄O₇·10 H₂O. It is found in some arid countries, as southern California and Tibet, but is now made commercially from the mineral colemanite, which is the calcium salt of a complex boric acid. When this is treated with a solution of sodium carbonate, calcium[Pg 266] carbonate is precipitated and borax crystallizes from the solution.

When heated borax at first swells up greatly, owing to the expulsion of the water of crystallization, and then melts to a clear glass. This glass has the property of easily dissolving many metallic oxides, and on this account borax is used as a flux in soldering, for the purpose of removing from the metallic surfaces to be soldered the film of oxide with which they are likely to be covered. These oxides often give a characteristic color to the clear borax glass, and borax beads are therefore often used in testing for the presence of metals, instead of the metaphosphoric acid bead already described.

Borax is extensively used as a constituent of enamels and glazes for both metal ware and pottery. It is also used as a flux in soldering and brazing, and in domestic ways it serves as a mild alkali, as a preservative for meats, and in a great variety of less important applications.

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 borate of silver is formed. If the solution of borax is dilute, however, an hydroxide of silver forms. Account for this difference in behavior.

[Pg 267]

CHAPTER XXII

THE METALS

The metals. The elements which remain to be considered are known collectively as the metals. They are also called the base-forming elements, since their hydroxides are bases. A metal may therefore be defined as an element whose hydroxide is a base. When a base dissolves in water the hydroxyl groups form the anions, while the metallic element forms the cations. From this standpoint a metal can be defined as an element 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 some conditions and as acids under others. We have seen that antimony is an element of this kind.

Occurrence of metals in nature. A few of the metals are found in nature in the free state. Among these are gold, platinum, and frequently copper. They are usually found combined with other elements in the form of oxides or salts of various acids. Silicates, carbonates, sulphides, and sulphates are the most abundant salts. All inorganic substances occurring in nature, whether they contain a metal or not, are called minerals. Those minerals from which a useful substance can be extracted are called ores of the substance. These two terms are most frequently used in connection with the metals.[Pg 268]

Extraction of metals,—metallurgy. The process of extracting a metal from its ores is called the metallurgy of the metal. The metallurgy of each metal presents peculiarities of its own, but there are several methods of general application which are very frequently employed.

1. Reduction of an oxide with carbon. Many of the metals occur in nature in the form of oxides. When these oxides are heated to a high temperature with carbon the oxygen combines with it and the metal is set free. Iron, for example, occurs largely in the form of the oxide Fe₂O₃. When this is heated with carbon the reaction expressed in the following equation takes place:

Many ores other than oxides may be changed into oxides which can then be reduced by carbon. The conversion of such ores into oxides is generally accomplished by heating, and this process is called roasting. Many carbonates and hydroxides decompose directly into the oxide on heating. Sulphides, on the other hand, must be heated in a current of air, the oxygen of the air entering into the reaction. The following equations will serve to illustrate these changes in the case of the ores of iron:

2. Reduction of an oxide with aluminium. Not all oxides, however, can be reduced by carbon. In such cases aluminium may be used. Thus chromium may be obtained in accordance with the following equation:

[Pg 269]

This method is a comparatively new one, having been brought into use by the German chemist Goldschmidt; hence it is sometimes called the Goldschmidt method.

3. Electrolysis. In recent years increasing use is being made of the electric current in the preparation of metals. In some cases the separation of the metal from its compounds is accomplished by passing the current through a solution of a suitable salt of the metal, the metal usually being deposited upon the cathode. In other cases the current is passed through a fused salt of the metal, the chloride being best adapted to this purpose.

Electro-chemical industries. Most of the electro-chemical industries of the country are carried on where water power is abundant, since this furnishes the cheapest means for the generation of electrical energy. Niagara Falls is the most important locality in this country for such industries, and many different electro-chemical products are manufactured there. Some industries depend upon electrolytic processes, while in others the electrical energy is used merely as a source of heat in electric furnaces.

Preparation of compounds of the metals. Since the compounds of the metals are so numerous and varied in character, there are many ways of preparing them. In many cases the properties of the substance to be prepared, or the material available for its preparation, suggest a rather unusual way. There are, however, a number of general principles which are constantly applied in the preparation of the compounds of the metals, and a clear understanding of them will save much time and effort in remembering the details in any given case. The most important of these general methods for the preparation of compounds are the following:[Pg 270]

1. By direct union of two elements. This is usually accomplished by heating the two elements together. Thus the sulphides, chlorides, and oxides of a metal can generally be obtained in this way. The following equations serve as examples of this method:

2. By the decomposition of a compound. This decomposition may be brought about either by heat alone or by the combined action of heat and a reducing agent. Thus when the nitrate of a metal is heated the oxide of the metal is usually obtained. Copper nitrate, for example, decomposes as follows:

Similarly the carbonates of the metals yield oxides, thus:

Most of the hydroxides form an oxide and water when heated:

When heated with carbon, sulphates are reduced to sulphides, thus:

3. Methods based on equilibrium in solution. In the preparation of compounds the first requisite is that the reactions chosen shall be of such a kind as will go on to completion. In the chapter on chemical equilibrium it was shown that reactions in solution may become complete in either of three ways: (1) a gas may be formed which escapes from solution; (2) an insoluble solid may be formed which precipitates; (3) two different ions may combine to form[Pg 271] undissociated molecules. By the judicious selection of materials these principles may be applied to the preparation of a great variety of compounds, and illustrations of such methods will very frequently be found in the subsequent pages.

4. By fusion methods. It sometimes happens that substances which are insoluble in water and in acids, and which cannot therefore be brought into double decomposition in the usual way, are soluble in other liquids, and when dissolved in them can be decomposed and converted into other desired compounds. Thus barium sulphate is not soluble in water, and sulphuric acid, being less volatile than most other acids, cannot easily be driven out from this salt When brought into contact with melted sodium carbonate, however, it dissolves in it, and since barium carbonate is insoluble in melted sodium carbonate, double decomposition takes place:

On dissolving the cooled mixture in water the sodium sulphate formed in the reaction, together with any excess of sodium carbonate which may be present, dissolves. The barium carbonate can then be filtered off and converted into any desired salt by the processes already described.

5. By the action of metals on salts of other metals. When a strip of zinc is placed in a solution of a copper salt the copper is precipitated and an equivalent quantity of zinc passes into solution:

In like manner copper will precipitate silver from its salts:

[Pg 272]

It is possible to tabulate the metals in such a way that any one of them in the table will precipitate any one following it from its salts. The following is a list of some of the commoner metals arranged in this way:

According to this table copper will precipitate bismuth, mercury, silver, or gold from their salts, and will in turn be precipitated by zinc, iron, tin, or lead. Advantage is taken of this principle in the purification of some of the metals, and occasionally in the preparation of metals and their compounds.

Important insoluble compounds. Since precipitates play so important a part in the reactions which substances undergo, as well as in the preparation of many chemical compounds, it is important to know what substances are insoluble. Knowing this, we can in many cases predict reactions under certain conditions, and are assisted in devising ways to prepare desired compounds. While there is no general rule which will enable one to foretell the solubility of any given compound, nevertheless a few general statements can be made which will be of much assistance.

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 mercurous chlorides. (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 only moderately soluble.)[Pg 273]

5. Sulphides. All sulphides are insoluble save those of ammonium, sodium, and potassium. The sulphides of calcium, barium, strontium, and magnesium are insoluble in water, but are changed by hydrolysis into acid sulphides which are soluble. On this account they cannot be prepared by precipitation.

6. Carbonates, phosphates, and silicates. All normal carbonates, phosphates, and silicates are insoluble save those of ammonium, sodium and potassium.

EXERCISES

1. Write equations representing four different ways for preparing Cu(NO₃)₂.

2. Write equations representing six different ways for preparing ZnSO₄.

3. Write equations for two reactions to illustrate each of the three ways in which reactions in solutions may become complete.

4. Give one or more methods for preparing each of the following compounds: CaCl₂, PbCl₂, BaSO₄, CaCO₃, (NH₄)₂S, Ag₂S, PbO, Cu(OH)₂ (for solubilities, see last paragraph of chapter). State in each case the general principle involved in the method of preparation chosen.

[Pg 274]

CHAPTER XXIII

THE ALKALI METALS

The family. The metals listed in the above table constitute the even family in Group I in the periodic arrangement of the elements, and therefore form a natural family. The name alkali metals is commonly applied to the family for the reason that the hydroxides of the most familiar members of the family, namely sodium and potassium, have long been called alkalis.

1. Occurrence. While none of these metals occur free in nature, their compounds are very widely distributed, being especially abundant in sea and mineral waters, in salt beds, and in many rocks. Only sodium and potassium occur in abundance, the others being rarely found in any considerable quantity.

2. Preparation. The metals are most conveniently prepared by the electrolysis of their fused hydroxides or chlorides, though it is possible to prepare them by reducing their oxides or carbonates with carbon.[Pg 275]

3. Properties. They are soft, light metals, having low melting points and small densities, as is indicated in the table. Their melting points vary inversely with their atomic weights, while their densities (sodium excepted) vary directly with these. The pure metals have a silvery luster but tarnish at once when exposed to the air, owing to the formation of a film of oxide upon the surface of the metal. They are therefore preserved in some liquid, such as coal oil, which contains no oxygen. Because of their strong affinity for oxygen they decompose water with great ease, forming hydroxides and liberating hydrogen in accordance with the equation

where M stands for any one of these metals. These hydroxides are white solids; they are readily soluble in water and possess very strong basic properties. These bases are nearly equal in strength, that is, they all dissociate in water to about the same extent.

4. Compounds. The alkali metals almost always act as univalent elements in the formation of compounds, the composition of which can be represented by such formulas as MH, MCl, MNO₃, M₂SO₄, M₃PO₄. These compounds, when dissolved in water, dissociate in such a way as to form simple, univalent metallic ions which are colorless. With the exception of lithium these metals form very few insoluble compounds, so that it is not often that precipitates containing them are obtained. Only sodium and potassium will be studied in detail, since the other metals of the family are of relatively small importance.

The compounds of sodium and potassium are so similar in properties that they can be used interchangeably for[Pg 276] most purposes. Other things being equal, the sodium compounds are prepared in preference to those of potassium, since they are cheaper. When a given sodium compound is deliquescent, or is so soluble that it is difficult to purify, the corresponding potassium compound is prepared in its stead, provided its properties are more desirable in these respects.

SODIUM

Occurrence in nature. Large deposits of sodium chloride have been found in various parts of the world, and the water of the ocean and of many lakes and springs contains notable quantities of it. The element also occurs as a constituent of many rocks and is therefore present in the soil formed by their disintegration. The mineral cryolite (Na₃AlF₆) is an important substance, and the nitrate, carbonate, and borate also occur in nature.

Preparation. In 1807 Sir Humphry Davy succeeded in preparing very small quantities of metallic sodium by the electrolysis of the fused hydroxide. On account of the cost of electrical energy it was for many years found more economical to prepare it by reducing the carbonate with carbon in accordance with the following equation:

The cost of generating the electric current has been diminished to such an extent, however, that it is now more economical to prepare sodium by Davy's original method, namely, by the electrolysis of the fused hydroxide or chloride. When the chloride is used the process is difficult to manage, owing to the higher temperature required to keep the electrolyte fused, and because of the corroding action of the fused chloride upon the containing vessel.

SIR HUMPHRY DAVY (English) (1778-1829)

Isolated sodium, lithium, potassium, barium, strontium, and calcium by means of electrolysis; demonstrated the elementary nature of chlorine; invented the safety lamp; discovered the stupefying effects of nitrous oxide

[Pg 277]

Fig. 76

Properties. Sodium is a silver-white metal about as heavy as water, and so soft that it can be molded easily by the fingers or pressed into wire. It is very active chemically, combining with most of the non-metallic elements, such as oxygen and chlorine, with great energy. It will often withdraw these elements from combination with other elements, and is thus able to decompose water and the oxides and chlorides of many metals.

Sodium peroxide (NaO). Since sodium is a univalent element we should expect it to form an oxide of the formula Na₂O. While such an oxide can be prepared, the peroxide (NaO) is much better known. It is a yellowish-white powder made by burning sodium in air. Its chief use is as an oxidizing agent. When heated with oxidizable substances it gives up a part of its oxygen, as shown in the equation

[Pg 278]

Water decomposes it in accordance with the equation

Acids act readily upon it, forming a sodium salt and hydrogen peroxide:

In these last two reactions the hydrogen dioxide formed may decompose into water and oxygen if the temperature is allowed to rise:

Sodium hydroxide (caustic soda) (NaOH). 1. Preparation. Sodium hydroxide is prepared commercially by several processes.

(a) In the older process, still in extensive use, sodium carbonate is treated with calcium hydroxide suspended in water. Calcium carbonate is precipitated according to the equation

The dilute solution of sodium hydroxide, filtered from the calcium carbonate, is evaporated to a paste and is then poured into molds to solidify. 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 reaction being expressed as follows:

[Pg 279]

The chlorine escapes as a gas, and by an ingenious mechanical device the sodium hydroxide is prevented from mixing with the salt in the solution.

In the Acker process the electrolyte is fused sodium chloride. The chlorine is evolved as a gas at the anode, while the sodium alloys with the melted lead which forms the cathode. When this alloy is treated with water the following reaction takes place:

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 into three compartments, A, B, and C. Compartment A is lined with magnesia brick. Compartments B and C are filled with melted lead, which also covers the bottom of A to a depth of about an inch. Above this layer in A is fused salt, into which dip carbon anodes D. The metallic box and melted lead is the cathode.

2. Properties. Sodium hydroxide is a white, crystalline, brittle substance 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. It is a strong base. It is used in a great many chemical industries, and under the name of lye is employed to a small extent as a cleansing agent for household purposes.

Sodium chloride (common salt) (NaCl). 1. Preparation. Sodium chloride, or common salt, is very widely distributed in nature. Thick strata, evidently deposited at one time by the evaporation of salt water, are found in many places. In the United States the most important localities for salt are New York, Michigan, Ohio, and Kansas. Sometimes the salt is mined, especially if it is in the pure form called rock salt. More frequently a strong brine is pumped from deep wells sunk into the salt deposit, and is then evaporated in large pans until the salt crystallizes out. The crystals are in the form of small cubes and contain no water of crystallization; some water is, however, held in cavities in the crystals and causes the salt to decrepitate when heated.

2. Uses. Since salt is so abundant in nature it forms the starting point in the preparation of all compounds[Pg 281] containing either sodium or chlorine. This includes many substances of the highest importance to civilization, such as soap, glass, hydrochloric acid, soda, and bleaching powder. Enormous quantities of salt are therefore produced each year. Small quantities are essential to the life of man and animals. Pure salt does not absorb moisture; the fact that ordinary salt becomes moist in air is not due to a property of the salt, but to impurities commonly occurring in it, especially calcium and magnesium chlorides.

Sodium sulphate (Glauber's salt) (Na₂SO₄·10H₂O). This salt is prepared by the action of sulphuric acid upon sodium chloride, hydrochloric acid being formed at the same time:

Some sodium sulphate is prepared by the reaction represented in the equation

The magnesium sulphate required for this reaction is obtained in large quantities in the manufacture of potassium chloride, and being of little value for any other purpose is used in this way. The reaction depends upon the fact that sodium sulphate is the least soluble of any of the four factors in the equation, and therefore crystallizes out when hot, saturated solutions of magnesium sulphate and sodium chloride are mixed together and the resulting mixture cooled.

Sodium sulphate forms large efflorescent crystals. The salt is extensively used in the manufacture of sodium carbonate and glass. Small quantities are used in medicine.

Sodium sulphite (Na₂SO₃·7H₂O). Sodium sulphite is prepared by the action of sulphur dioxide upon solutions[Pg 282] of sodium hydroxide, the reaction being analogous to the action of carbon dioxide upon sodium hydroxide. Like the carbonate, the sulphite is readily decomposed by acids:

Because of this reaction sodium sulphite is used as a convenient source of sulphur dioxide. It is also used as a disinfectant and a preservative.

Sodium thiosulphate (hyposulphite of soda or "hypo") (Na₂S₂O₃·5H₂O). This salt, commonly called sodium hyposulphite, or merely hypo, is made by boiling a solution of sodium sulphite with sulphur:

It is used in photography and in the bleaching industry, to absorb the excess of chlorine which is left upon the bleached fabrics.

Sodium carbonate (sal soda)(Na₂CO₃·10H₂O). There are two different methods now employed in the manufacture of this important substance.

1. Le Blanc process. This older process involves several distinct reactions, as shown in the following equations.

(a) Sodium chloride is first converted into sodium sulphate:

[Pg 283]

(b) The sodium sulphate is next reduced to sulphide by heating it with carbon:

(c) The sodium sulphide is then heated with calcium carbonate, when double decomposition takes place:

2. Solvay process. This more modern process depends upon the reactions represented in the equations

The reason the first reaction takes place is that sodium hydrogen carbonate is sparingly soluble in water, while the other compounds are freely soluble. When strong solutions of sodium chloride and of ammonium hydrogen carbonate are brought together the sparingly soluble sodium hydrogen carbonate is precipitated. This is converted into the normal carbonate by heating, the reaction being represented in the second equation.[Pg 284]

[Pg 285]

Properties of sodium carbonate. Sodium carbonate forms large crystals of the formula Na₂CO₃ · 10 H₂O. It has a mild alkaline reaction and is used for laundry purposes under the name of washing soda. Mere mention of the fact that it is used in the manufacture of glass, soap, and many chemical reagents will indicate its importance in the industries. It is one of the few soluble carbonates.

Sodium hydrogen carbonate (bicarbonate of soda) (NaHCO₃). This salt, commonly called bicarbonate of soda, or baking soda, is made by the Solvay process, as explained above, or by passing carbon dioxide into strong solutions of sodium carbonate:

The bicarbonate, being sparingly soluble, crystallizes out. A mixture of the bicarbonate with some substance (the compound known as cream of tartar is generally used) which slowly reacts with it, liberating carbon dioxide, is used largely in baking. The carbon dioxide generated forces its way through the dough, thus making it porous and light.

Sodium nitrate (Chili saltpeter) (NaNO₃). This substance is found in nature in arid regions in a number of places, where it has been formed apparently by the decay of organic substances in the presence of air and sodium salts. The largest deposits are in Chili, and most of the nitrate of commerce comes from that country. Smaller deposits occur in California and Nevada. The commercial salt is prepared by dissolving the crude nitrate in water,[Pg 286] allowing the insoluble earthy materials to settle, and evaporating the clear solution so obtained to crystallization. The soluble impurities remain for the most part in the mother liquors.

Since this salt is the only nitrate found extensively in nature, it is the material from which other nitrates as well as nitric acid are prepared. It is used in enormous quantities in the manufacture of sulphuric acid and potassium nitrate, and as a fertilizer.

Sodium phosphate (Na₂HPO₄·12H₂O). Since phosphoric acid has three replaceable hydrogen atoms, three sodium phosphates are possible,—two acid salts and one normal. All three can be made without difficulty, but disodium phosphate is the only one which is largely used, and is the salt which is commonly called sodium phosphate. It is made by the action of phosphoric acid on sodium carbonate:

It is interesting as being one of the few phosphates which are soluble in water, and is the salt commonly used when a soluble phosphate is needed.

Normal sodium phosphate (Na₃PO₄). Although this is a normal salt its solution has a strongly alkaline reaction. This is due to the fact that the salt hydrolyzes in solution into sodium hydroxide and disodium phosphate, as represented in the equation

Sodium hydroxide is strongly alkaline, while disodium phosphate is nearly neutral in reaction. The solution as a whole is therefore alkaline. The salt is prepared by adding a large excess of sodium hydroxide to a solution of disodium[Pg 287] phosphate and evaporating to crystallization. The excess of the sodium hydroxide reverses the reaction of hydrolysis and the normal salt crystallizes out.

Sodium tetraborate (borax) (Na₂B₄O₇·10H₂O). The properties of this important compound have been discussed under the head of boron.

POTASSIUM

Occurrence in nature. Potassium is a constituent of many common rocks and minerals, and is therefore a rather abundant element, though not so abundant as sodium. Feldspar, which occurs both by itself and as a constituent of granite, contains considerable potassium. The element is a constituent of all clay and of mica and also occurs in very large deposits at Stassfurt, Germany, in the form of the chloride and sulphate, associated with compounds of sodium and magnesium. In small quantities it is found as nitrate and in many other forms.

The natural decomposition of rocks containing potassium gives rise to various compounds of the element in all fertile soils. Its soluble compounds are absorbed by growing plants and built up into complex vegetable substances; when these are burned the potassium remains in the ash in the form of the carbonate. Crude carbonate obtained from wood ashes was formerly the chief source of potassium compounds; they are now mostly prepared from the salts of the Stassfurt deposits.

Preparation and properties. The metal is prepared by the same method used in the preparation of sodium. In most respects it is very similar to sodium, the chief difference being that it is even more energetic in its action upon other substances. The freshly cut, bright surface instantly becomes dim through oxidation by the air. It decomposes water very vigorously, the heat of reaction being sufficient to ignite the hydrogen evolved. It is somewhat lighter than sodium and is preserved under gasoline.

Fig. 78

Potassium hydroxide (caustic potash) (KOH). Potassium hydroxide is prepared by methods exactly similar to those[Pg 289] used in the preparation of sodium hydroxide, which compound it closely resembles in both physical and chemical properties. It is not used to any very great extent, being replaced by the cheaper sodium hydroxide.

Action of the halogen elements on potassium hydroxide. When any one of the three halogen elements—chlorine, bromine, and iodine—is added to a solution of potassium hydroxide a reaction takes place, the nature of which depends upon the conditions of the experiment. Thus, when chlorine is passed into a cold dilute solution of potassium hydroxide the reaction expressed by the following equation takes place:

If the solution of hydroxide is concentrated and hot, on the other hand, the potassium hypochlorite formed according to equation (1) breaks down as fast as formed:

Equation (1), after being multiplied by 3, may be combined with equation (2), giving the following:

This represents in a single equation the action of chlorine on hot, concentrated solutions of potassium hydroxide. By means of these reactions one can prepare potassium chloride, potassium hypochlorite, and potassium chlorate. By substituting bromine or iodine for chlorine the corresponding compounds of these elements are obtained. Some of these compounds can be obtained in cheaper ways.

If the halogen element is added to a solution of sodium hydroxide or calcium hydroxide, the reaction which takes place is exactly similar to that which takes place with[Pg 290] potassium hydroxide. It is possible, therefore, to prepare in this way the sodium and calcium compounds corresponding to the potassium compounds given above.

Potassium chloride (KCl). This salt occurs in nature in sea water, in the mineral sylvine, and, combined with magnesium chloride, as carnallite (KCl·MgCl₂·6H₂O). It is prepared from carnallite by saturating boiling water with the mineral and allowing the solution to cool. The mineral decomposes while in solution, and the potassium chloride crystallizes out on cooling, while the very soluble magnesium chloride remains in solution. The salt is very similar to sodium chloride both in physical and chemical properties. It is used in the preparation of nearly all other potassium salts, and, together with potassium sulphate, is used as a fertilizer.

Potassium bromide (KBr). When bromine is added to a hot concentrated solution of potassium hydroxide there is formed a mixture of potassium bromide and potassium bromate in accordance with the reactions already discussed. There is no special use for the bromate, so the solution is evaporated to dryness, and the residue, consisting of a mixture of the bromate and bromide, is strongly heated. This changes the bromate to bromide, as follows:

The bromide is then crystallized from water, forming large colorless crystals. It is used in medicine and in photography.

Potassium iodide (KI). Potassium iodide may be made by exactly the same method as has just been described for the bromide, substituting iodine for bromine. It is more frequently made as follows. Iron filings are [Pg 291]treated with iodine, forming the compound Fe₃I₈; on boiling this substance with potassium carbonate the reaction represented in the following equation occurs:

Potassium iodide finds its chief use in medicine.

Potassium chlorate (KClO₃). This salt, as has just been explained, can be made by the action of chlorine on strong potassium hydroxide solutions. The chief use of potassium chlorate is as an oxidizing agent in the manufacture of matches, fireworks, and explosives; it is also used in the preparation of oxygen and in medicine.

Potassium nitrate (saltpeter) (KNO₃). This salt was formerly made by allowing animal refuse to decompose in the open air in the presence of wood ashes or earthy materials containing potassium. Under these conditions the nitrogen in the organic matter is in part converted into potassium nitrate, which was obtained by extracting the mass with water and evaporating to crystallization. This crude and[Pg 292] slow process is now almost entirely replaced by a manufacturing process in which the potassium salt is made from Chili saltpeter:

This process has been made possible by the discovery of the Chili niter beds and the potassium chloride of the Stassfurt deposits.

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 of gunpowder.

Potassium cyanide (KCN). When animal matter containing nitrogen is heated with iron and potassium carbonate, complicated changes occur which result in the formation of a substance commonly called yellow prussiate of potash, which has the formula K₄FeC₆N₆. When this substance is heated with potassium, potassium cyanide is formed:

Since sodium is much cheaper than potassium it is often used in place of it:

The mixture of cyanides so resulting serves most of the purposes of the pure salt. It is used very extensively in several metallurgical processes, particularly in the extraction of gold. Potassium cyanide is a white solid characterized by its poisonous properties, and must be used with extreme caution.

Potassium carbonate (potash) (K₂CO₃). This compound occurs in wood ashes in small quantities. It cannot be prepared by the Solvay process, since the acid carbonate is quite soluble in water, but is made by the Le Blanc process. Its chief use is in the manufacture of other potassium salts.[Pg 294]

Other salts of potassium. Among the other salts of potassium frequently met with are the sulphate (K₂SO₄), the acid carbonate (KHCO₃), the acid sulphate (KHSO₄), and the acid sulphite (KHSO₃). These are all white solids.

LITHIUM, RUBIDIUM, CÆSIUM

Of the three remaining elements of the family—lithium, rubidium, and cæsium—lithium is by far the most common, the other two being very rare. Lithium chloride and carbonate are not infrequently found in natural mineral waters, and as these substances are supposed to increase the medicinal value of the water, they are very often added to artificial mineral waters in small quantities.

COMPOUNDS OF AMMONIUM

General. As explained in a previous chapter, when ammonia is passed into water the two compounds combine to form the base NH₄OH, known as ammonium hydroxide. When this base is neutralized with acids there are formed the corresponding salts, known as the ammonium salts. Since the ammonium group is univalent, ammonium salts resemble those of the alkali metals in formulas; they also resemble the latter salts very much in their chemical properties, and may be conveniently described in connection with them. Among the ammonium salts the chloride, sulphate, carbonate, and sulphide are the most familiar.

Ammonium chloride (sal ammoniac) (NH₄Cl). This substance is obtained by neutralizing ammonium hydroxide with hydrochloric acid. It is a colorless substance crystallizing in fine needles, and, like most ammonium salts, is very soluble in water. When placed in a tube and heated strongly it decomposes into hydrochloric acid and ammonia. When these gases reach a cooler portion of the tube they[Pg 295] at once recombine, and the resulting ammonium chloride is deposited on the sides of the tube. In this way the salt can be separated from nonvolatile impurities. Ammonium chloride is sometimes used in preparation of ammonia; it is also used in making dry batteries and in the laboratory as a chemical reagent.

Ammonium sulphate ((NH₄)₂SO₄). This salt resembles the chloride very closely, and, being cheaper, is used in place of it when possible. It is used in large quantity as a fertilizer, the nitrogen which it contains being a very valuable food for plants.

Ammonium carbonate ((NH₄)₂CO₃). This salt, as well as the acid carbonate (NH₄HCO₃), is used as a chemical reagent. They are colorless solids, freely soluble in water. The normal carbonate is made by heating ammonium chloride with powdered limestone (calcium carbonate), the ammonium carbonate being obtained as a sublimate in compact hard masses:

The salt always smells of ammonia, since it slowly decomposes, as shown in the equation

The acid carbonate, or bicarbonate, is prepared by saturating a solution of ammonium hydroxide with carbon dioxide:

It is a well-crystallized stable substance.

Ammonium sulphide ((NH₄)₂S). Ammonium sulphide is prepared by the action of hydrosulphuric acid upon ammonium hydroxide:

[Pg 296]

If the action is allowed to continue until no more hydrosulphuric acid is absorbed, the product is the acid sulphide, sometimes called the hydrosulphide:

If equal amounts of ammonium hydroxide and ammonium acid sulphide are brought together, the normal sulphide is formed:

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 is much used in the process of chemical analysis because it is a soluble sulphide and easily prepared. On exposure to the air ammonium sulphide slowly decomposes, being converted into ammonia, water, and sulphur:

As fast as the sulphur is liberated it combines with the unchanged sulphide to form several different ammonium sulphides in which there are from two to five sulphur atoms in the molecule, thus: (NH₄)₂S₂, (NH₄)₂S₃, (NH₄)₂S₅. These sulphides in turn decompose by further action of oxygen, so that the final products of the reaction are those given in the equation. A solution of these compounds is yellow and is sometimes called yellow ammonium sulphide.

Fig. 79

[Pg 298]

The spectra of the known elements have been carefully studied, and any element which imparts a characteristic color to a flame, or has a spectrum of its own, can be identified even when other elements are present. Through the spectroscopic examination of certain minerals a number of elements have been discovered by the observation of lines which did not belong to any known element. A study of the substance then brought to light the new element. Rubidium and cæsium were discovered in this way, rubidium having bright red lines and cæsium a very intense blue line. Lithium colors the flame deep red, and has a bright red line in 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 brought about, giving the general principle involved in each change: NaCl --> Na₂SO₃, Na₂SO₃ --> NaCl, NaCl --> NaBr, Na₂SO₄ --> NaNO₃, NaNO₃ --> NaHCO₃.

3. What carbonates are soluble?

4. State the conditions under which the reaction represented by the following equation can be made to go in either direction:

5. Account for the fact that solutions of sodium carbonate and potassium carbonate are alkaline.

6. What non-metallic element is obtained from the deposits of Chili saltpeter?

7. Supposing concentrated hydrochloric acid (den. = 1.2) to be worth six cents a pound, what is the value of the acid generated in the preparation of 1 ton of sodium carbonate by the Le Blanc process?

8. What weight of sodium carbonate crystals will 1 kg. of the anhydrous salt yield?

9. Write equations for the preparation of potassium hydroxide by three different methods.

10. What would take place if a bit of potassium hydroxide were left exposed to the air?

11. Write the equations for the reactions between sodium hydroxide and bromine; between potassium hydroxide and iodine.

12. Write equations for the preparation of potassium sulphate; of potassium acid carbonate.

ROBERT WILHELM BUNSEN (German) (1811-1899)

Invented many lecture-room and laboratory appliances (Bunsen burner); invented the spectroscope and with it discovered rubidium and cæsium; greatly perfected methods of electrolysis, inventing a new battery; made many investigations among metallic and organic substances

[Pg 299]

13. What weight of carnallite would be necessary in the preparation of 1 ton of potassium carbonate?

14. Write the equations showing how ammonium chloride, ammonium sulphate, ammonium carbonate, and ammonium nitrate may be prepared from ammonium hydroxide.

15. Write an equation to represent the reaction involved in the preparation of ammonia from ammonium chloride.

16. What substances already studied are prepared from the following compounds? ammonium chloride; ammonium nitrate; ammonium nitrite; sodium nitrate; sodium chloride.

17. How could you prove that the water in crystals of common salt is not water of crystallization?

18. How could you distinguish between potassium chloride and potassium iodide? between sodium chloride and ammonium chloride? between sodium nitrate and potassium nitrate?

[Pg 300]

CHAPTER XXIV

THE ALKALINE-EARTH FAMILY

The family. The alkaline-earth family consists of the very abundant element calcium and the much rarer elements strontium and barium. They are called the alkaline-earth metals because their properties are between those of the alkali metals and the earth metals. The earth metals will be discussed in a later chapter. The family is also frequently called the calcium family.

1. Occurrence. These elements do not occur free in nature. Their most abundant compounds are the carbonates and sulphates; calcium also occurs in large quantities as the phosphate and silicate.

2. Preparation. The metals were first prepared by Davy in 1808 by electrolysis. This method has again come into use in recent years. Strontium and barium have as yet been obtained only in small quantities and in the impure state, and many of their physical properties,[Pg 301] such as their densities and melting points, are therefore imperfectly known.

3. Properties. The three metals resemble each other very closely. They are silvery-white in color and are about as hard as lead. Their densities increase with their atomic weights, as is shown in the table on opposite page. Like the alkali metals they have a strong affinity for oxygen, tarnishing in the air through oxidation. They decompose water at ordinary temperatures, forming hydroxides and liberating hydrogen. When ignited in the air they burn with brilliancy, forming oxides of the general formula MO. These oxides readily combine with water, according to the equation

Each of the elements has a characteristic spectrum, and the presence of the metals can easily be detected by the spectroscope.

4. Compounds. The elements are divalent in almost all of their compounds, and these compounds in solution give simple, divalent, colorless ions. The corresponding salts of the three elements are very similar to each other and show a regular variation in properties in passing from calcium to strontium and from strontium to barium. This is seen in the solubility of the sulphate and hydroxide, and in the ease of decomposition of the carbonates, as given in the table. Unlike the alkali metals, their normal carbonates and phosphates are insoluble in water.

CALCIUM

Occurrence. The compounds of calcium are very abundant in nature, so that the total amount of calcium in the earth's crust is very large. A great many different compounds[Pg 302] containing the clement are known, the most important of which are the following:

Preparation. Calcium is now prepared by the electrolysis of the melted chloride, the metal depositing in solid condition on the cathode. It is a gray metal, considerably heavier and harder than sodium. It acts upon water, forming calcium hydroxide and hydrogen, but the action does not evolve sufficient heat to melt the metal. It promises to become a useful substance, though no commercial applications for it have as yet been found.

Calcium oxide (lime, quicklime) (CaO). Lime is prepared by strongly heating calcium carbonate (limestone) in large furnaces called kilns:

When pure, lime is a white amorphous substance. Heated intensely, as in the oxyhydrogen flame, it gives a brilliant light called the lime light. Although it is a very difficultly fusible substance, yet in the electric furnace it can be made to melt and even boil. Water acts upon lime with the evolution of a great deal of heat,—hence the name quicklime, or live lime,—the process being called slaking. The equation is

Lime readily absorbs moisture from the air, and is used to dry moist gases, especially ammonia, which cannot be[Pg 303] dried by the usual desiccating agents. It also absorbs carbon dioxide, forming the carbonate

CaO + CO₂ = CaCO₃.