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Practical chemistry

Newell, Lyman Churchill, 1867-1933

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chemistry

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PRACTICAL CHEMISTRY

BY LYMAN C. NEWELL, Ph.D. (johns hopkins)

PROFESSOR OF CHEMISTRY, BOSTON UJaVERSITV, BOSTON, MASS.

AUTHOR OF

"EXPERIMENTAL CHEXHSTRY,"' "DESCRIPTIVE CHEMISTRY"

" GENERAL CHEMISTRY," " INORGANIC CHEMISTRY FOR

COLLEGES," "LABORATORY MANUAL OF

INORGANIC CHEMISTRY"

D. C. HEATH & CO., PUBLLSHERS

BOSTON Xt.W YORK CHICAGO

^ ( 33

PREFACE

A TEXTBOOK intended for the first year of chemistry should nu-it certain requirements. Both pupil and teacher must be considin-d. The pupil needs a book which is written plainly, illustrated fully, and applied practically; which stimulates interest in chemistrv, and which can be studied profitably. The teacher needs a b;)ok in which the text is carefully selected, judiciously apportioned, and properlv arranged ; which provides teaching material, and which can be taught successfully. The author believes his Practical Chemistry meets these requirements.

The text includes the facts and principles suitable for beirinncrs. No important topic has been omitted. The selection is varied. i)er- mitting the use of the book in different kinds of classes. The topics suggested by the College Entrance Examination Board and the Board of Regents (New York) have been incorporated.

The style and arrangement are clear — short sentences, brief para- graphs, explanatory examples, lettered subdivisions, numbered sec- tions, instructive topical headings. Unusual care has been taken to write clearly such topics as symbols, formulas, equations, atomic weights, molecular weights, and valence.

The exercises, which are placed at the end of the chapters, have l)ecn prepared to meet the real needs of pupils. The author believes that a class should be drilled on fundamentals as well as provided with opportunities to answer test questions. Hence, numerous exercises have been provided for review, home study, practical e.xpcrience. and themes. Teachers are urged to examine these exercises and as- sign those best adapted to the needs of the class. The list of books incorporated in the Appendix, § 5, will be serviceable in connection with home study and other assignments.

The problems emphasize fundamental principles and processes; many require original thinking. The liberal selection will meet the requirements of different kinds of classes.

An examination of the book will show that it is a Practical Chem- istry— practical in several ways. The more vital applications of chemistry are described in connection with the appropriate facts or principles. IMoreover, the major applications of chemistry are so fully described and illustrated that a pupil can obtain from the book itself an accurate and adequate account of modern chemical indus-

iv PREFACE

tries. Teachers are referred to such topics as purification of water (69), nitric acid (196), sulphur (250), sulphuric acid (268, 270), fuel oil and gasolene (311, 312. Fig. 112), sodium silicate (387), glass (394), cast iron (490), steel (494. 495) ^ and copper (528, Fig. 193). ' Certain features of this book should not be overlooked.

(i) The study of carbon is begun in Chapter III, thereby introduc- ing a famihar element and its striking applications at the start. This element is reviewed (with extensions) in Chapter XXI, and its useful compounds are studied in Chapters XXII (Fuels and Illuminants), XXIII (Other Carbon Compounds), and XXIV (Food). No apol- ogy is offered for devoting so much space to carbon. It is a funda- mental element, and the practical apphcations of this element and its compounds are indispensable and conspicuous.

(2) The portion of the text that deals with theory is thoroughly adequate. The material is strategically distributed among several chapters (VII, VIII, X, XV, XVI, XVII). This distribution, as well as the simple treatment, will make the theory of chemistry less formidable to beginners.

(3) Certain metals are treated rather fully and placed in the first part of the text devoted to metals. These are sodium, calcium, iron, aluminium, and copper (Chapters XXIX-XXXIII) . By this ar- rangement the essential principles of metallurgy, the characteristics of metals, and the apphcations of metals and their compounds are available for study before the hurried work at the close of the year.

(4) Half-tone illustrations have been liberally used because they tell a story which is often more complete than verbal description. The author trusts they will be made indispenable adjuncts to instruc- tion. The portraits of chemists, which were made from originals in the author's private collection, have been inserted to arouse interest in those chemists who have contributed conspicuously to the founding and progress of chemistry.

(5) The drawings of processes and apparatus are likewise intended as a supplement to the text. Especial care has been taken to draw them accurately and to represent salient features without confusing details. These drawings, with minor exceptions, were made by Har- old C. Spencer, Boston University, 1914- The author is deeply in- debted to him for commendable patience and skillful cooperation.

The author is grateful for suggestions and assistance received from members of the faculty of Boston University and from former stu- dents who are teaching chemistry in high schools.

L. C. N.

Boston, Mass. May, 1922.

ACKNO\M.EDGMEXTS

The author acknowledges the courtesy and Kcnerosily of the fol- lowing for the use of photographs : -- Acheson Graphite Co., Niagara Falls, X. V. ( loo) ; Air Reduction Sales Co., N. Y. (28, 125) ;

American Agricultural Chemical Co., Boston, Mass. (152, i53) ; American La France Fire Engine Company of Massachusetts, Boston,

Mass. (18); Board of Water Supply — City of New York, X. \ . (34) ; Bureau of Mines, Washington, D. C. (iij ; Commonwealth Steel Co., St. Louis, Mo. (181); Coors Porcelain Co., Golden, Colo. (192) ;

Electro Bleaching Gas Co., N. Y. (63) ;

Freeport Sulphur Co., N. Y. (89) ;

Honolulu Iron Works, N. Y. (132, i33. i34) ;

Illinois Glass Co., and Editor of Bottles, Alton, 111. (146) ;

Libbev-Owens Sheet Glass Co.. Charleston, W . \ a. (145) ;

Lurav Caverns Corporation, Luray, Va. (162) ;

Macbeth, Evans Co., Pittsburgh, Pa. (144) "' . ^ ^ ,

Ma-nesia Association of America, Philadelphia, Pa. (200, 201);

Metal & Thermit Co., N. Y. (189 — left, 190) ;

National Lime Association, Washington, D. C. (160) ;

New England Oil Refining Co., Boston, Mass. (iii, 113);

Niagara Alkali Co., Niagara Falls, N. Y. (157) ; , ^ . . .

Oxygen Products Co. and Electrolabs Co., Pittsburgh, Pa. (4),

Petroleum Age, Chicago, III. (no) ;

Philadelphia Quartz Co., Philadelphia, Pa. (143);

Raritan Copper Works, Perth Amboy, N. J. (194, i95, ^^^^ ^°^' °'\-°'^

Rhode Island Agricultural Experiment Station, Kingston, R. I. (lOi),

Riter-Conley Co., Pittsburgh, Pa. (172) ;

Semet-Solvay Co., Syracuse, N. Y. (104) ;

Swift & Co., Chicago, 111. (138) ;

Texas Gulf Sulphur Co., N. Y. (90) ;

Thermal Syndicate, Ltd., N. Y. (142) ; - u* a\-

Underwood & Underwood, N. Y. (183, 184, 185 copyrighted) .

Vulcan Iron Works, Wilkes-Barre, Pa. (167) ;

Wallace & Tiernan Co., Inc., Newark. N. J. ^3:^) .

S. D. Warren Co., Boston, Mass. (136, i37) ;

Welsbach Co., Gloucester, N. J. (131);

AiIa Wood Iron & steel Co., Philadelphm I.-.. (..^V.

Youngstown Sheet & Tube Co., YounRstown, Oh.o (■„, .,S. ,8.),

CH.\PTER I.

II. III.

IV.

VI.

VII.

VIII.

IX.

XI.

XII.

XIII.

XIV.

XV.

XVI.

XVII.

XVIII.

XIX.

XX.

XXI.

XXII. XXIII.

CONTENTS

Chemical Ch.\nge — Elements — Compoun'ds —

Law of Constant Composition . . . .

Oxygen

Carbon and Its Oxides — Law of Multiple Pkd-

portions Hydrogen .... Measurement of Gases . Water — Law of Gas X'olumes — Hydrogen

Peroxide

Atomic Theory — Atoms and Molecules

Symbols axd Formulas ....

Xitrogen — Air — Argon — Helium — Liquid Air

Reactions and Equations

Chlorint: — Hydrochloric Acid

Acids — Bases — Salts — Neutralization

Ammonia — Ammonium Compounds .

Nitric Acid — Nitrogen Oxides .

Molecular Weights and For.milas

Atomic Weights

Valence ....

Ions and Ionization

Sulphur — Sulphides

Sulphur Oxides, Acids, and Salts

Carbon — Carbonic Acid - Carbon atks

Oxides — Carbides

Fuels — Illuminants — Fla.mk

Other Carbon Compounds . . . •

^5 50

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1 1 >

I :')

I V)

1 7 ' 10:

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CONTENTS

CHAPTER

XXIV.

XXV.

XXVL

XXMI. XXVIII.

XXIX.

XXX.

XXXI.

XXXII.

XXXIII.

XXXIV.

XXXV.

XXXVI.

XXXVII.

XXXVIII.

Food

Silica — Silicates — Glass

Metals and Non-Metals — Periodic Classi- fication .......

Fluorine — Bromine — Iodine .

Phosphorus — Phosphates — Arsenic — Anti MONY — Bismuth . . .

Sodium — Potassium ......

Calcium Carbonate — Lime — Cement .

Iron and Steel

Aluminium — Clay — Porcelain

Copper

Magnesium — Zinc — INIercury

Tin — Lead

SiL\'ER — Photography — Gold

Chromium — IVIanganese — Platinum - Nickel — Cobalt

Radium and Radioactivity

Appendix

Metric System

Thermometers

Weights of Gases

\"apor Pressure

Books . . . • • • •

Index

PAGE 316

337

346 360 376 391 415 426

439 452 464

474 482

493 493 493 494 494 497

PRACTICAL CHEMISTRY

CHAPTER I

CHEMICAL CHANGE — COMPOUNDS — ELEMENTS

1. What is chemistry? — If we light a candle and let it burn, the candle wax slowly disappears — '' burns away.'' Let us study this change carefully. By holding a cold, dry bottle over the burning candle (Fig. i), a him of water gathers on the inside. Now if we remove the bottle, pour in some clear hmewater, and shake the bottle, the clear hquid becomes cloudy. This cloudiness is caused by a gas called carbon dioxide. Hence, the candle wax did not really " burn away." By burning in the air, the wax changed into water and carbon dioxide.

The experiment just studied is an example of chemical change. Chem- istry is the science devoted largely to a study of chemical changes.

2. Substances and properties. — The dilYercnt mate- rials involved in a chemical change are called substances. We are able to recognize and distinguish substances by characteristics called properties. Some properties readily detected, e.g. color, taste, odor, hardnes^ physical state {i.e. if solid, hquid

Fig.

Studying a chemical change

and

are the

gaseous )

Other

PRACTICAL CHEMISTRY

properties are found by measurement or special experiment, and are usually expressed by numbers, e.g. boiling point, melting point, conductivity of heat or electricity, and spe- cific gravity (i.e. relative weight). These properties, often called physical properties, are important, and we often use them in describing and recognizing substances.

Very important, also, are chemical properties — the char- acteristics of substances revealed in chemical changes. Thus, iron rusts in the air but gold does not, copper dis- solves readily in nitric acid but platinum does not. Chemi- cal properties are the different items, so to speak, in the

chemical conduct of substances. In chem- istry we pay special at- tention to chemical con- duct.

3. What is the first characteristic of chemi- cal change ? — Let us answer this question by an experiment. Iron and sulphur are readily recognized. Besides their conspicuous prop- erties, sulphur dis- solves in carbon disulphide, and iron is attracted by a magnet. If we heat a mixture of iron powder and sulphur in a test tube (Fig. 2), the mass soon glows, and the glow often spreads and continues after the tube is removed from the flame. If we break the test tube and examine the product, we find that it is a hard, black solid, quite unlike either iron or sulphur. If we grind it in a mortar, and add a few drops of acid to a little of the

Fig. 2. — Heating a mixture of iron and sulphur in a test tube

CHEMICAL CHANGE ^COMPOUNDS — ELEMENTS 3

powder, a gas with a bad odor is liberated. No such gas is produced when acid is added to the iron, the sulphur, or the mixture. Moreover, this black solid is not attracted by a magnet nor does it dissolve in carbon disulphide.

Clearly a new substance has been formed. We may de- scribe the chemical change in chemical language. Two different substances, iron and sulphur, each with char- acteristic properties, have disappeared and a new substance called iron sulphide has been produced with properties different from the properties of the original substances. All chemical changes, however complicated, involve trans- formation of substances. So we may say : —

Tlie first characteristic of chemical change is the forma- tion of one or more new substances from the original substance or substances.

4. What is the second characteristic of chemical change ? — The experiment just considered may be done with more care. If we take the right proportions of iron and sulphur (7 of iron and 4 of sulphur), weigh each substance accu- rately, and heat the mixture until the chemical change is complete, we find that the weight of the iron sulphide is the same as the weight of the iron and sulphur together. Precise study has shown that this striking observation is true of all chemical changes. Hence,

The second characteristic of chemical change is unchanged weight.

5. An important law. — We see, then, as a result of our study, that in a chemical change substances are trans- formed but this transformation involves no change in the total weight of all the substances involved. This general fact about chemical change is often stated as the law of the conservation of matter : —

No weight is lost or gained in a chemical change.

4 PRACTICAL CHEMISTRY

6. What is a substance ? — The word substance is often used in ordinary language to designate various kinds of materials. Thus, wood, cloth, paper, flour, soil, sugar, iron, and water are substances. But in chemistry the word substance means a kind of material which is alike through- out, i.e. all parts have the same properties. Moreover all specimens of a given substance have the same set of proper- ties. Sulphur is a substance. That is, every particle of sulphur is Uke every other particle, and different samples of sulphur are alike. Anything found in a specimen of sulphur that has different properties is another substance. It is called an impurity and the specimen of sulphur is called im- pure sulphur.

7. What is a mixture? — ^ If we examine certain ma- terials, we find that they consist of two or more sub- stances, which can be easily recognized with the eye (or a magnifying glass) or can be separated by sifting, dis- solving, or filtering. A material which consists of sub- stances merely mixed or stuck together is called a mixture. Paint, milk, soil, granite, and muddy water are examples of mixtures.

Mixtures have certain unmistakable characteristics. First, the ingredients of a mixture may vary in kind and proportion. Thus, soil may be largely sand, or clay, or organic matter. Indeed, the kind and proportion of in- gredients may vary widely and the mixture still have the same name. Second, the ingredients can be readily sepa- rated. For example, the preparation of flour consists mainly in separating the starch from the other parts of the ground kernel. Third, when we compare a mixture with a (chemical) substance, we find a marked difference. Every specimen and each part of a substance has the same proper- ties. But the properties of a mixture vary with the kind

CHEMICAL CHANGE — COMPOUNDS — ELEMENTS 5

and proportion of the ingredients. Obviously, a mixture of iron and sulphur may consist of any proportion of the ingredients, and the properties will depend on the proportion ; whereas iron sulphide is always alike through- out.

8. Compounds and elements. — There are two classes of substances — compounds and elements. How are they related? If we heat strongly some red powder, called mercury oxide, in a test tube (Fig. 2), the powder turns dark and soon minute silvery drops collect on the upper part of the test tube. Now if we push well down into the test tube a joss stick with a glowing end, the glow increases and the end of the joss stick bursts into a flame. Clearly a gas is being produced which differs from ordinary air.

The new substances formed from the red powder are the liquid, mercury, which can be seen, and the colorless gas, oxygen, which mingles with the air. Both substances are quite different from the original red solid, mercury oxide. Therefore, by heating we have transformed the substance, mercury oxide, into two different substances, mercury and oxygen. A chemical change has taken place. Mercury oxide has been decomposed into mercury and oxygen. Moreover, these new substances, mercury and oxygen, differ from mercury oxide in one fundamental property, viz. they cannot be decomposed.

The relation of mercury and oxygen to mercury oxide is clear. Mercury and oxygen are the fundamental constit- uents of mercury oxide. We can decompose mercury oxide into its constituents, but there the decomposition stops. We have reached the end, chemically speaking. These two substances, mercury and oxygen, are not only simpler than mercury oxide, but they are examples of the simplest sub-

6 PRACTICAL CHEMISTRY

stances. Substances which have not yet been decom- posed by chemists into simpler substances are called ele- ments, whereas substances which can be decomposed into elements are called compounds. Obviously the compound mercury oxide consists of the elements mercury and oxy- gen.

9. What are the characteristics of compounds? — Com- pounds have several essential characteristics.

(i) The elements in compounds are chemically united. That is, the constituents of a compound are not merely mingled or lying side by side as in a mixture. They are held together by a force which is sometimes called chemical attraction. To separate a compound into its elem.ents, this attraction must be reduced or overcome chemically. Only by a chemical change can we decompose a compound into its elements or unite elements into a compound. Whereas a mixture can usually be separated into its in- gredients by some simple mechanical operation (7) .

(2) The properties of compounds differ — often con- spicuously — from the properties of the elements that compose them. Thus, the blue sohd copper sulphate is composed of three elements — the red metal copper, the yellow sohd sulphur, and the colorless gas oxygen.

(3) The most important characteristic of chemical com- pounds is their constant composition. This means that any given chemical compound always consists of an un- varying per cent of the same elements. For example, the compound water always contains 88.82 per cent of the element oxygen and 11. 18 per cent of the element hydrogen. This is an important fact in chemistry, so important that it is stated as the law of constant composition : —

A chemical compound has a constant cofuposition by weight. This law will be studied again (98).

CHEMICAL CHANGE — COMPOUNDS — ELEMENTS 7

10. There is a large number of compounds. — The num- ber of compounds formed by various combinations of the elements is very large indeed, — too large for one person to study profitably. Fortunately, we can learn much about the essentials of chemistry by confining our study to the important elements and the best representatives of a few classes of compounds.

11. More about elements. — There is only a small number of elements — about eighty-five. A complete table of the elements will be found on the inside of the back cover of this book.

Important elements are shown in the accompanying table.

T.\BLE OF Important Elements and Their Symbols

Element

Symbol

Element

Symbol

Element

Symbol

Aluminium

Gold

Oxygen

Barium

Hydrogen

Phosphorus

Bromine

Iodine

Potassium

Calcium

Iron

Silicon

Carbon

Lead

Silver

Chlorine

Magnesium

Sodium

Chromium

:Manganese

Sulphur

Copper

Mercury

Tin

Fluorine

Nitrogen

Zinc

Less than half of the elements are common. About 98 per cent of the outer shell of the earth's crust consists of compounds derived from eight elements. The per cent of these elements is shown in the accompanying table.

Table of the Composition of the Earth's Crust

Oxygen 47.07

Silicon 28.06

Aluminium 7.90

Iron

Calcium . . . Potassium .

4-43 3-44 2.45

Sodium 2.43

Magnesium 2.40

Remainder 1.82

PRACTICAL CHEMISTRY

Compounds in the ocean likewise consist of but few ele- ments, as appears from the following : —

Table of the Approximate Composition of the Ocean

Oxygen 85.79

Hydrogen 10.67

Chlorine 2.07

Sodium 1. 14

^Magnesium ... .0.14 Sulphur 0.09

Calcium 0.05

Bromine 0.008

Carbon 0.002

The human body is a complicated structure, and yet the compounds in it are composed of but few elements; and some of these elements are in very small proportions. This fact is apparent from the illustrative table.

Table of the Average Composition of the Human Body

Oxygen 65.00

Carbon 18.00

Hydrogen 10.00

Nitrogen 3.00

Calcium 2 .00

Phosphorus .... i.oo

Potassium 0.35

Sulphur 0.25

Sodium 0.15

Chlorine o. i ^

Magnesium 0.05

Iron 0.004

Iodine trace

Fluorine trace

Silicon trace

It is evident from these tables that only a dozen elements are abundant. As a matter of fact a few elements by their various combinations furnish most of the substances studied in this book. The elements given in the above tables, in- deed most elements, do not occur free. That is, they do not occur singly but are combined with other elements in the form of compounds. Only a very few elements are found in the free, native, or uncombined state, e.g. sulphur, copper, gold, and carbon in the earth's crust, and oxygen and nitrogen in the atmosphere.

12. Each element has a symbol. — Each element is desig- nated by an abbreviation called a symbol (11). Symbols are usually the first letter, or the first and a significant one, of the common name of the element. Thus, O is the symbol

CHEMICAL CHANGE— COMPOUNDS -ELEMENTS 9

of oxygen, C of carbon, CI of chlorine, Zn of zinc. In some cases the symbol is an abbreviation of a Latin name, e.g. Cu for copper {cuprum), Fe for iron ifcrrum), A^ for silver {argentum). All symbols, as we shall see later (105), have a very explicit meaning. They are used constantly, not only to designate elements and compounds l)ut also to express the relations of elements and compounds in chemical change. The symbol of each important element should be learned early in the study of chemistry. (See Exercise 11 at the end of this chapter.)

13. Each compound has a formula. — Just as each element is designated by a symbol, so each compound is represented by a formula: A formula is a group of sym- bols — the symbols of the elements of which the compound is composed. Thus, FeS is the formula of the compound iron sulphide, HgO of mercuric oxide, CO2 of carbon dioxide, and H2O of water. Later we shall learn how formulas are derived.

14. How chemical change, elements, and compounds are related. — In the chemical change involving iron and sulphur (3), these elements unite to form the compound iron sulphide. So also in the chemical change illustrated by the behavior of mercuric oxide (8), this compound decomposes into the elements mercury and oxygen. We cannot have a chemical change without involving one or more compounds. And in many cases elements, too, arc directly concerned.

Chemical change is sometimes called chemical action. The term reaction, or interaction, is usually applied to a single chemical change. A reaction may be rei)resented in a condensed form, thus : —

Sulphur + Iron = Iron Sulphide Such a condensed expression is called an equation. In

lO PRACTICAL CHEMISTRY

these simple equations, the plus sign may also be read and and the equality sign form {s). Thus, sulphur and iron form iron sulphide. Or in full, the elements iron and sulphur under suitable conditions undergo a chemical change which results in the formation of the compound iron sulphide.

Since no weight is lost or gained in a chemical change (5), an equation can be used to represent this funda- mental fact. Experiment shows that 32 parts of sulphur always unite with 56 parts of iron. This fact is expressed as an equation thus : —

Sulphur + Iron = Iron Sulphide 32 56 88

This equation is read; 32 parts of sulphur and 56 parts of iron form (or equal) 88 parts of iron sulphide. By parts we mean any denomination by weight, e. g. grams, kilo- grams, pounds, tons.

EXERCISES

1. State characteristic properties of (a) glass (&) gasolene, (c) water, {d) paper, (e) air, (/) lead.

2. Define the term substance as used in chemistry.

3. Give three illustrations of chemical change you have observed.

4. State and illustrate the two characteristics of chemical change.

5. State the law of the conservation of matter. What does the word conservation mean ?

6. Name (a) five elements and {h) five compounds.

7. What are some characteristics of a mJxture? Name three familiar mixtures.

8. How can water be distinguished from gasolene? Copper from iron? Glass from sand? Air from illuminating gas? Sugar from starch?

9. Name the eight most abundant elements in the earth's crust in their order. Give the symbol of each.

10. What is an element? A compound? In what fundamental way do elements and compounds differ? Could you prepare (a) a compound from elements, {h) elements from a compound, (c) com- pounds from compounds, {d) elements from elements?

CHEMICAL CHANGE — COMPOUNDS— ELEMENTS ii

11. Learn the symbol of each element mentioned in this chapter. What is the formula of iron sulphide, water, mercury oxide, carbon diox- ide ?

12. What is a reaction? An equation? Give an example of each.

13. Interpret this equation : Mercury Oxide = Mercury + Oxygen.

14. Make a list of all the new words in this chapter and defme each.

15. State practical uses of chemistry in (a) the school building, (b) your home, (r) agriculture, (d) manufacturing processes in your city or town.

PROBLEMS

{The Metric System of Weights and Measures is constantly used in Chemistry and it should he learned or reviewed at once. See Appendix, §i.)

1. What is the abbreviation of gram, centigram, liter, meter, cubic centimeter, decimeter, milligram, millimeter?

2. Express (a) i liter in cubic centimeters, {b) 2 1. in cc, (c) i meter in centimeters, (d) 250 cm. in dm., (e) 1 kg. in grams, (/) 250 gm. in mg., (g) 56.75 1. in cc, (//) 1250 cc. in I., (/) i cc. in cu. m.

3. How many cc. in (a) i liter, (b) 1 cu. dm., (r) i cu. m. ?

4. If I m. of magnesium ribbon weighs 4 dg., how many mg. will 5 cm. weigh?

5. Into how many pieces 5 cm. long can a glass tube i m. long be cut?

6. A flask holds 750 cc. Express its capacity in (a) 1., (b) cu. dm.

7. A bottle holds exactly 1250 cc. (a) How many grams of water will fill it? (b) How many kg.? (c) How many 1.?

8. Suppose exactly 3.5 gm. of iron and 2 gm. of sulphur are heated until the chem.ical change is complete. What weight of iron sulphide is produced? (Suggestion. See 5.)

9. Suppose 5 gm. of iron and 2 gm. of sulphur are heated until the chemical change is complete, (a) What weight of iron sulphide is pro- duced? (b) Is any sulphur or iron left over? (c) If so, which and how much?

10. What per cent of iron sulphide is sulphur? Iron?

11. When mercury oxide is made, 25 parts by weight of mercury and 2 of oxygen unite. What is the per cent of (a) mercury and (b) oxy- gen in mercury oxide?

12. How many grams of (a) hydrogen and (b) oxygen can be ob- tained from 150 gm. of water? (Suggestion. See 9 (3).)

CHAPTER II

OXYGEN

15. Oxygen is an abundant and important element. — It forms nearly 21 per cent (by volume) of the atmosphere. Combined with hydrogen, it constitutes 88.82 per cent (by weight) of water ; combined with silicon and certain

metals, it makes up nearly half of the earth's crust (11). Combined with car- bon, hydrogen, nitrogen, and other elements, it forms a large part of ani- mal and vegetable matter. Thus, the human body contains about 65 per cent of combined oxygen, while vegetable matter contains about 40 per cent.

If the elements were to be named in order of importance, oxygen would be first. Without free

oxygen animal life is impossible. It is also necessary in

the famihar process of burning.

16. Preparation of oxygen. — Among the first to pre- pare oxygen w^as the English scientist, Priestley (Fig. 3). He prepared it in 1774 by heating mercury oxide in much the same way as previously described (8).

Fig. 3. — Priestley (i 733-1804)

OXYGEN 13

The gas can be prepared l)y decomposing other compounds of oxygen, such as potassium chlorate (KCIO3), lead oxide (PbOa), or barium oxide (BaO..). Thus, potassium chlorate — a compound of oxygen, chlorine, and potassium — when heated to a moderately high temperature yields all its oxygen, and a compound, potassium chloride (KCl), remains.

Oxygen can also be prepared from water (HoO). When an electric current is passed through water which con- tains sulphuric acid or sodium hydroxide, two gases, oxygen

Fig. 4. — A plant for manufacturing oxygen (and hydrogen) by passing an electric current through water containing sodium hydroxide

and hydrogen, are liberated in separate tubes or compart- ments. This method is used to prepare oxygen (and hydro- gen) on a large scale (Fig. 4). When a small quantity is needed, it is prepared by dropping water upon sodium peroxide.

17. Preparation of oxygen in the laboratory. — Oxygen is conveniently prepared in the laboratory by heating a mixture of potassium chlorate and manganese dioxide in the apparatus shown in Fig. 5. A mixture of about four

PRACTICAL CHEMISTRY

parts of potassium chlorate and one part of manganese dioxide is put in the test tube A and gently heated.

The oxygen escapes through the delivery tube D into bottles previously filled with water and inverted over the end of the tube in the pneumatic trough. The oxygen bubbles up into the

Apparatus for preparing oxygen ^^^^^^ ^^^ displaces in the laboratory

the water.

18. The preparation of oxygen illustrates chemical change. — The chemical change consists in the decompo- sition of the compound potassium chlorate into the element oxygen, and the compound potassium chloride. This chemical change may be expressed by this equation : —

Potassium Chlorate = Oxygen + Potassium Chloride

(Potassium-Chlorine-Oxygen) (Potassium-Chlorine)

Chemical changes like this are common, and the term decomposition is applied to them. Decomposition is chemical change in which a compound is separated chemi- cally into other substances which are elements or com- pounds.

19. Properties of oxygen. — Pure oxygen has no color, odor, or taste ; certain impurities may give a slight odor and taste to the gas prepared in the laboratory. It is not very soluble in water, and for this reason can be collected over water. Oxygen is slightly heavier than air. One liter of oxy- gen weighs 1.43 grams, if carefully measured and weighed at the temperature of o degrees as registered by a centigrade

OXYGEN

thermometer and also under a pressure of 760 millimeters as registered by a barometer (or briefly at 0° C. and 760 mm.). This value — 1.43 grams — should be remembered. 20. The chemical conduct of oxygen is conspicuous. — Oxygen forms compounds with most elements. It also in- teracts chemically with many compounds. This combining or interacting is often strik- ing on account of the accompanying Hght and heat. At ordinary temperatures oxy- gen unites slowly with many elements. Thus, metals, such as lead, zinc, and cop- per, tarnish or rust slowly, i.e. they com- bine slowly with the oxygen of the air. With phosphorus, however, the chemical action is quite rapid, as may be seen by p^^ 6.^^^1phur the glow and fumes when the end of a burning in a phosphorus- tipped match is rubbed, espe- bottle of oxygen cially in the dark.

The chemical activity of oxygen at temperatures above the ordi- nary is readily shown by putting burning or glowing substances into the gas. The action, scarcely noticeable in air, becomes energetic, and the substances burn rapidly and brilliantly. Thus, a glowing splint of wood or piece of charcoal when put into a bottle of oxygen bursts immediately into flame. Sulphur burns in air with a feeble bluish flame, but in oxygen the Clt^ flame becomes large and brilliant (Fig. 6). Iron

/ \ can hardly be made to burn in air, but if steel wool

(matted strands of iron) is merely heated and thrust into a bottle of oxygen, the iron burns, sends off a shower of sparks, and often forms drops of molten iron which crack the bottle (Fig. 7).

Fig. 7. — Iron burning in a bottle of oxy- gen

In the experiments just described, we should note especially that the oxygen itself

i6 PRACTICAL CHEMISTRY

does not burn. Rather it assists burning. So if we were to state briefly the chemical conduct of oxygen, we would say oxygen does not burn but assists the burning of other substances. The chemical conduct of oxygen described above is sometimes called its chief chemical property.

21. Test for oxygen. — The critical examination made to estabhsh the identity of an element or compound is called testing or making a test. The behavior of a sub- stance under stated conditions is called the test for the sub- stance. Thus, the test for oxygen is its chief chemical con- duct, viz. the gas does not burn, but assists burning at elevated temperatures.

22. The chief chemical property of oxygen illustrates chemical change. — In the experiments described in 20, one feature is conspicuous, viz. the disappearance of the original substances and the formation of new substances. The chemical change in the case of the carbon, sulphur, and iron is the combining of oxygen with these elements. The oxygen unites with each element, and the product is a com- pound of the two elements. This kind of chemical change illustrates combination, and can be expressed by an equa- tion. Thus : —

Carbon -j- Oxygen = Carbon Dioxide

(Carbon-Oxygen)

Combination is chemical change in which compounds are formed by the union of two or more substances.

23. Oxidation and oxides. — The special terrn oxida- tion is applied to those cases of combination in which oxygen is the combining element. Substances which furnish the oxygen are oxidizing agents. Free oxygen and air are oxidizing agents, though the oxygen for oxidation is often provided by compounds of oxygen, especially those that yield oxygen readily, such as potassium chlorate. The

OXYGEN 17

compound formed by the union of oxygen and another ele- ment is called an oxide of that element. Thus, carbon forms carbon dioxide.

Oxides of different elements are distinguished by placing the name of the element (or a slight modification of it) before the word oxide, e.g. magnesium oxide, nitric oxide. Sometimes di-, or a similar numerical syllable, is prefixed to the word oxide, e.g. manganese dioxide (MnO^), sul- phur trioxide (SO3), phosphorus pentoxide (P2O5). Many familiar substances are oxides, e.g. water is hydrogen oxide (H2O), lime is calcium oxide (CaO), sand is siUcon dioxide (SiO^X red lead is lead tetroxide (Pb304), and the most abundant iron ore is iron oxide (FcoOs).

24. Oxidation of compounds. — Oxidation is by no means limited to elements. Many compounds burn readily, i.e. combine as a whole with oxygen. Thus, carbon monox- ide (CO) unites directly with oxygen to form carbon diox- ide (CO2).

Oxidation is an important chemical change. Later we shall find that the term includes other kinds of chemical action.

25. Combustion is oxidation. — During oxidation heat is liberated, and if the heat is intense, hght is also produced. Different substances react with oxygen, i.e. oxidize, at dif- ferent rates. If oxidation is slow, as in the rusting of some metals, the temperature may not rise appreciably, because the heat escapes about as fast as it is liberated. If oxida- tion is rapid, heat is liberated quickly, the temperature rises suddenly, and the substance burns, often with dazzhng light. This rapid oxidation that produces heat and Ught is called combustion. In ordinary language combustion means fire or burning ; in chemical language it is rapid oxidation.

PRACTICAL CHEMISTRY

Sometimes the heat Hberated during slow oxidation cannot es- cape readily, but accumulates, hastens the oxidation, and finally the temperature rises to such a point that the substance takes fire. Thus, oily rags carelessly thrown aside by painters, moist hay stored in a poorly ventilated barn, and soft coal kept in a pile a long time in the air or in the w^arm hold of a ship sometimes take fire without apparent cause. Such fires, often unexpected and disastrous, are said to be due to spontaneous combustion, though they are simply cases of slow oxidation which becomes accelerated by accumulated heat.

26. Combustion was first interpreted by Lavoisier.

— Simple as it is, the answer to the question " What hap- pens when a substance burns? " was not made quickly. Indeed, ' several famous chemists tried to answer this fascinating question and their com- bined work extended over about a century. The answer was delayed many years by a false theory called the phlogiston theory. The advocates of this theory believed that '' combustible substances contain a principle called phlogiston, and that when a substance burns, phlogiston escapes." This is false. But it was not until about 1775 that the French chemist' Lavoi- sier (Fig. 8) proved by his own and others' experiments: (i) that phlogiston did not exist, and (2) that ordinary combustion is a process of combining with " a certain sub- stance contained in the air." Soon after, he showed that this *' substance " is identical with the gas discovered by

Fig. 8. — Lavoisier (i 743-1 794)

OXYGEN

Fig. 9. — Apparatus used by La- voisier in his famous experiment on combustion

Priestley in 1774. In 1778 Lavoisier named the gas oxygen.

So important is Lavoisier's discovery we must consider it in detail.

Previous to Lavoisier's time it had been shown that air is necessary for combustion. It had also been shown that when a substance burns in air, the product weighs more than the original substance. Lavoisier himself verified this second observation by care- ful experiments. He knew, then, these two simple facts about combustion: d) air is necessary and (2) the weight increases. Still, the question to be answered was, '' What happens when a sub- stance burns?" Lavoisier answered the question by a conclusive experiment (Fig. 9).

He put mercury in the retort having a long neck which communi- cated with the jar. The jar contained air confined over mercury in the larger lower vessel. Having arranged the apparatus so that the mercury in the retort was in contact with the air in both the retort and jar, he heated the retort by means of the furnace and kept the mer- cury in the retort just below the boiling point for twelve days. When the apparatus was cold, he noticed two things : (i) A red powder had accumulated on the surface of the mercury in the retort, and (2) the air in the retort and jar had decreased in volume about one fifth.

He collected the red powder, put it in a glass vessel, heated it intensely, collected the gas, and measured its volume. He found the volume was the same as the decrease in the volume of the original air inclosed in the apparatus. He tested this gas, and found it had the properties of the gas previously obtained by Priestley, especially the property of making a lighted candle burn vigorously. Also he tested the gas left behind in the retort and jar and found that it ex-

20 PRACTICAL CHEMISTRY

tinguished a lighted candle. From these facts Lavoisier concluded : (i) that air contains a gas which is removed by heating substances in it, (2) that this gas is about one fifth of air, (3) that this gas unites with substances in the process called combustion or burning, and (4) that this gas is identical with the gas obtained by Priestley from his " red precipitate of mercury."

We might summarize Lavoisier's work thus : oxygen is the gas in the air that is necessary for combustion.

27. Oxygen is essential to life. — Free ox>^gen is essen- tial to all forms of animal Hfe. If human beings or animals are deprived of air, they die. In breathing, air is drawn into our lungs ; here the oxygen of the air is taken up by the blood, which distributes it to all parts of the body. This oxygen slowly oxidizes the tissues of the body. By this slow oxidation, waste products are formed and heat is suppHed to the body. Two of these waste products are carbon dioxide gas and water vapor, which are exhaled from the lungs ; water vapor is also given off through the skin. New tissue is built up from the food we eat.

The human body resembles a steam engine. In each, the oxygen of the air helps burn fuel largely composed of carbon. In the engine, the products escape through a chimney, and the heat produced by the chemical change is used to form steam, which moves parts of the machine. In the body, the products escape through the lungs and other organs, and the heat keeps the body at the temperature at which it can best perform its functions.

28. Manufacture and uses of oxygen. — Oxygen for industrial and scientific use is manufactured by decompos- ing water by an electric current (16) or by separating air into its components (128). The gas is stored under pres- sure in strong metal tanks (Fig. 10).

A mixture of oxygen and hydrogen or acetylene, if burned in a suitable apparatus, produces an intensely hot flame. The oxy-hydrogen flame is used to melt certain

OXYGEN

Fig. lo. — A tank of oxygen

metals and to produce the intense light of the stereopticon, while the oxy-acetylene flame finds appUcation in welding and in burning apart heavy steel structures, e.g. girders of bridges (57, 334).

Oxygen is often administered to persons who are too ill or weak to inhale the ordi- nary volume of air. In submarine boats the oxygen of the air used up is replaced by oxygen released from tanks. Airplanes are often equipped with an oxygen tank and breathing apparatus to supply quickly the oxygen needed by the aviator at eleva- tions where the air is rarefied.

29. Saving lives by oxygen apparatus. — Oxygen is used in various forms of breath- ing apparatus for rescue work. The pulmotor, or lung- motor, is essentially a pump by which air rich in oxygen can be forced into the lungs at intervals approximating the normal rate of breathing. The pulmotor is used to resuscitate persons who have been overcome by smoke or poisonous gases (e.g. illuminating gas) or who have been rendered unconscious by drowning or by an electric shock. Fire departments, poKce officials, and public health officers are suppUed with pulmotors for emergencies.

Another form of rescue apparatus can be hung from the shoulders like a knapsack. The man's head in one type is covered with a leather helmet ; in a more recent type the man's nose is clipped so he must breathe through his mouth (Fig. ii). Flexible tubes connect his mouth with a breathing bag (right), a cyUnder of compressed oxygen gas and a regenerating can (left) ; the latter contains po- tassium hydroxide to absorb the water vapor and carbon dioxide exhaled from the lungs. The amount of oxygen

22 PRACTICAL CHEMISTRY

needed by the wearer (as well as the circulation of the gases) can be regulated by a valve on the cylinder ; the nitrogen originally inhaled and in the apparatus at first is

^ «r^^^P^

k'^i

■i^

■L^ M

^^®

Fig. II. — Man equipped %Yith oxygen-breathing apparatus. He breathes through his mouth from the gas bag carried on his chest (right). The tank of oxygen and the regenerating can are carried on his back (left)

breathed over and over. The cyHnder contains enough oxygen for about two hours.

A man provided with an oxygen-breathing apparatus can safely enter places where the air contains smoke or poisonous gas, and make repairs, extinguish fires, or res- cue workmen who have been overcome. Extensive use is made of this kind of rescue apparatus in mine disasters.

30. Ozone. — Ozone is a gas related to oxygen. It is formed from oxygen when electric sparks pass through air, e.g. during a thunder storm. It is also formed near electrical machinery, and is prepared on a commercial scale by a silent electrical discharge, i.e. in a special kind of machine which permits a discharge without sparks.

Ozone changes back into oxygen — slowly at ordinary temperatures and rapidly at high temperatures. Hence ozone is a good oxidizing agent. It is sometimes used to bleach {i.e. whiten) flour, starch, oils,

OXYGEN 23

waxes, and wool. Its main use is to destroy bacteria in drinking water, especially in certain European cities.

EXERCISES

1. Prepare a brief summary of this chapter.

2. What is the characteristic chemical property of oxygen?

3. If air contains a large proportion of another gas besides oxygen, how must the general properties of this other ingredient compare with those of oxygen?

4. Make a list of the new chemical words in this chapter and de- fine each.

5. (a) Make a list of the name and symbol of each element men- tioned in studying oxygen, (b) Make a list of the compounds men- tioned in this chapter and the elements in each, (c) Learn the for- mula of each compound (as given).

6. What general chemical change is involved in ordinary burning? What class of chemical changes is illustrated by (a) preparation of oxy- gen from mercuric oxide, (b) burning of sulphur in oxygen?

7. Cite cases of spontaneous combustion of which you have heard or read. Suggest methods to prevent spontaneous combustion of (c) oily rags, (b) coal, (c) hay.

8. Devise an experiment to show that air contains oxygen.

9. Topics for home study, (a) Why does a draft of air make a fire burn \vell? (b) A draft of air often extinguishes a candle flame. Why? (c) What oxides are found in the home? (d) How do fish obtain oxygen ?

10. What chemical part does oxygen take in : (a) respiration,

(b) burning, (c) combustion, (d) oxidation, (e) kindling a fire?

11. Define (a) oxidation and (b) oxide. Name five oxides.

12. Essay topics : (a) Uses of oxygen, (b) Discovery of oxygen.

PROBLEMS

1. (a) What is the weight in gm. of 35 1. of oxygen? (6) Of 35,000 cc? (c) Of 35 cubic decimeters?

2. How many gm. of oxygen are in a bottle holding 2.5 1. (at 0° C. and 760 mm.) ?

3. A pupil prepared enough oxygen to fill a tank holding i cu. m. (at 0° C. and 760 mm.). How many gm. were prepared?

4. (a) How many liters (at 0° C. and 760 mm.) will 25 gm. of oxygen occupy? (b) How many gm. will 25 1. of oxygen weigh?

24 PRACTICAL CHEMISTRY

5. How many gm. of oxygen (at o° C. and 760 mm.) in a cylindri- cal gas holder which is i m. high and 30 cm. in diameter?

6. If air contains 21 per cent of oxygen by volume, how many gm. of oxygen can be extracted from 950 1. of air?

7. A pupil prepared five bottles of oxygen, each holding 250 cc. (at 0° C. and 760 mm.). How many gm. of oxygen were prepared?

8. Water contains 88.82 per cent of oxygen. Suppose 0.5 kg. was decomposed, how many liters of oxygen (at 0° C. and 760 mm.) were formed?

. 9. A room is 10 m. long, 5 m. wide, and 4 m. high. How many gm. of water must be decomposed to furnish enough oxygen (at 0° C. and 760 mm.) to fill the room?

10. Potassium, chlorate contains 39.18 per cent of oxygen. If 35 gm. are heated, (a) how many gm. of oxygen are liberated, and (b) how many bottles each containing 250 cc. will the gas fill?

CHAPTER III CARBON AND ITS OXIDES

31. Carbon is an important and useful element. — Like oxygen, the element carbon is found both free and com- bined. Free carbon is famiUar as the black soUd that makes up the greater part of the common substances coal, char- coal coke, and lampblack. The soft, slippery, shiny solid called graphite is also largely carbon. And. strange as it may seem, the valuable gem called diamond is also carbon — pure crystalline carbon.

Coal, charcoal, and coke are fuels. Lampblack is made into printing ink and black paint. The various varieties of these three forms of carbon are used in many chemical industries, e.g. the manufacture of iron and steel. Graphite is made into the 'Mead" of lead pencils, stove polish, lubri- cants, crucibles (for melting metals), and electrodes (for electric furnaces and electrolytic cells (497, 512)). Impure diamond is used to poKsh pure diamond and other gems, to cut glass, and as the cutting part in the diamond drill that is used to bore wells for oil and water and to drill hard rocks.

32. Carbon forms an extraordinarily large number of compounds. — The natural and manufactured compounds of carbon number over 200,000. These compounds are so numerous they are treated by themselves in a branch of chemistry called Organic Chemistry. Carbon is the fundamental element of plant and animal life. Compounds composed of carbon, hydrogen, and oxygen, and in some

26 PRACTICAL CHEMISTRY

cases also of nitrogen, form such common substances as sugar, starch, fat, cotton, paper, flour, rubber, soap, wool, and meat. With hydrogen it forms a large class of com- pounds called hydrocarbons, which are found in illumi- nating and fuel gases, petroleum, kerosene, gasolene, lubri- cating oils, paraffin wax, and turpentine. The manufac- tured compounds of carbon include dyes, medicines, per- fumes, and a vast number of other substances.

The commonest inorganic compounds — those usually considered in the branch of chemistry we are studying — are the carbonates and the oxides. The carbonates are compounds of carbon, oxygen, and a metal such as calcium, magnesium, or sodium. Thus, calcium carbonate (CaCOs) is the natural substance called limestone, marble, or chalk. Sodium carbonate (NasCOs) is the common substance washing soda. There are two carbon oxides — carbon diox- ide (CO2) and carbon monoxide (CO).

In later chapters we shall consider more fully, carbon and some of its important compounds. In this chapter we shall Umit our study to carbon and its two oxides.

33. Formation of carbon dioxide. — We have already seen that carbon dioxide is formed when carbon burns in oxygen and also when a candle burns in air (20, 1). These are examples of a common chemical change, viz. the for- mation of the compound carbon dioxide when the element carbon, or a combustible compound of carbon, burns in oxygen or in air.

We have also found that the process called burning or combustion consists usually in the union of carbon (and also, of course, hydrogen, if present) with oxygen (25, 26). This means that carbon dioxide is being constantly formed by the burning of such common fuels as wood, paper, coke, coal, charcoal, oil, and gas. In fact, carbon dioxide is

CARBON AND ITS OXIDES

always one of the products of combustion, as they are often called, yielded by burning any substance which contains carbon, e.g. sugar, starch, wax, meat, milk, camphor, alcohol, oil, dyes, fat, and drugs. The presence of carbon dioxide can be shown by bubbling the products of com- bustion, e.g. smoke, through Umewater (1).

Reference was made in 27 to the digestion of food. One of the products of this complex process is carbon dioxide. This means in simple language that the element carbon, which is one of the con- stituents of our food, is oxidized to carbon dioxide. Heat is liberated when carbon is oxidized, and this chemical change is one part of the continuous and complicated process by which our bodies keep uni- formly warm. The two main products cf digestion are exhaled from^ the lungs. The presence of carbon dioxide in exhaled breath may be readily shown by blowing the breath gently through a glass tube into a bottle containing limewater (Fig. 12). The Uquid becomes milky, owing to the reaction between the carbon dioxide and the Umewater. Limewater contains cal- cium hydroxide, which reacts with the carbon dioxide and forms the white insoluble compound calcium carbonate. This chemical change is a test for carbon dioxide (from any source). We express the chemical change thus : —

Carbon Dioxide -h Calcium Hydroxide = Calcium Carbonate-f Water

(Carbon-Oxygen) (Calcium-Hydrogen- (Calcium-Carbon-Oxypcn) (Hydro-

Oxygen) gen-Oxygen)

Carbon dioxide is also formed by other chemical changes, such as the decay of many kinds of animal and vegetable

Fig. 12. — Blow- ing through limewater to show the pres- ence of carbon dioxide in the breath

PRACTICAL CHEMISTRY

matter and fermentation of organic substances like sugar. The latter process is illustrated by the liberation of carbon dioxide during the raising of bread.

34. Preparation of carbon dioxide. — We can prepare carbon dioxide by burning carbon, or a combustible com- pound of carbon, in oxygen or air, but the gas could not be conveniently separated from the smoke or the other gases, such as the unused oxygen or the nitrogen (in the air). On a large scale, however, the gas is prepared by passing air over hot coke, purifying the gaseous product, and finally extracting the carbon dioxide by a special cooling process.

In the laboratory carbon dioxide is most conveniently prepared by the interaction of an acid and a carbonate. Dilute hydrochloric

acid and calcium carbonate (in the form of marble chips) are usually used. When the acid is poured on the calcium car- bonate, the gas is rapidly lib- erated. The apparatus shown in Fig. 13 can be used for this experiment. Calcium carbonate is put in the bottle A and the dilute hydrochloric acid is in- troduced through the dropping

tube B by pressing the clamp.

Fig. 13. — Apparatus for preparing 1. i- • 1 4.

u A- -A f -A ^r.A The carbon dioxide passes out

carbon dioxide from acid and ^

marble in the laboratory through the delivery tube D

into the pneumatic trough, bub- bles up into the bottles (previously filled with water), and displaces the water. The bottles when full of gas are immediately removed and covered with filter paper or a glass plate.

35. Some experiments with carbon dioxide. — If several bottles of carbon dioxide are collected, we can perform experiments which will give us information about the properties of this gas.

(i) A blazing joss stick or splinter of wood, if plunged into a bottle of carbon dioxide, is immediately extinguished. (2) If we lower a short, lighted candle into a bottle of air and quickly invert a bottle of

CARBON AND ITS OXIDES

carbon dioxide loosely over it, the gas falls down upon the candle which is soon extinguished. (3) When we fill a bottle of carbon diox ide one third full of water, cover it tightly with the hand, shake vigo rously a minute or two, then invert the bottle in a dish of water, and remove the hand, water rushes up into the bottle.

36. Some properties of carbon dioxide.

— Carbon dioxide is a colorless and

odorless gas. It is heavier than air —

about 1.5 times heavier. A hter of the

pure gas weighs 1.98 grams (at 0° C. and

760 mm.). It dissolves in water. At

ordinary temperature and pressure, water

dissolves about its own volume of carbon

dioxide. Under increased pressure the

solubihty increases, but the gas escapes

when the pressure is lessened. Soda water is manufactured

by dissolving carbon dioxide under pressure in water ;

hence when soda water is drawn from a soda fountain or siphon (Fig. 14), the water bubbles and forms a froth, owing to the escape of the gas under diminished pres- sure. Many beverages, such as ginger ale, are "carbonated," i.e. they are manufactured by forcing carbon dioxide into the prepared liquid ; the bottle is closed tightly with a cork or cap. When the bottle is opened, the gas which

Fig. 14. — A siphon of soda water

escapes (Fig. 15) will turn lime-

Fig. 15. — E.xperiment to show that carbon dioxide escapes from a carbonated water milky ; hence the gas is beverage carbon dioxide.

30 PRACTIC\L CHEMISTRY

37. Liquid and solid carbon dioxide. — Unlike oxygen, carbon diox- ide can be readily liquefied and solidified. If enough pressure is applied at ordinary temperatures, the gas becomes a liquid. Liquid carbon dioxide is stored and sold in strong steel tanks. If the tank is properly opened, part of the escaping liquid by evaporating quickly removes so much heat that the remainder becomes white, snowlike, soHd carbon dioxide.

38. Chemical conduct of carbon dioxide. — In our ex- periments we noticed that carbon dioxide, unlike oxygen, did not assist combustion. Nor did it burn. This nega- tive behavior, so to speak, is sometimes called inertness. We may say, then, that carbon dioxide is an inert gas. However, it does react with some substances. Thus, it combines with water to form a compound called carbonic acid ; but this compound is not stable, i.e. it decomposes readily and re-forms the water and carbon dioxide.

We have called attention several times to the reaction in which carbon dioxide and Hmewater form calcium carbonate and water. This chemical change, we have also said, serves as a test for carbon dioxide. A similar reaction takes place between carbon dioxide and sodium hydroxide. The prod- uct, in this case, however, is sodium carbonate. Sodium carbonate is soluble in water and the solution feels slippery, like soap ; in fact, sodium carbonate is sometimes called washing soda and is used in large quantities as a cleansing agent. Sodium bicarbonate is closely related to sodium carbonate. Sodium bicarbonate is cooking soda. Alone, or as an ingredient of baking powder, it is widely used in cooking because it gives off carbon dioxide which puffs up the dough.

39. Relation of carbon dioxide to life. — Carbon diox- ide is not poisonous, though the presence of a small quan- tity in the air of a room is objectionable. As already stated, the carbon dioxide that is exhaled from our lungs

CARBON AND ITS OXIDES

is one of the products formed by the oxidation of the tis- sues of the body, new tissue itseh* being formed from the food (27). The carbon needed for the rebuilding of tissue is supphed by starch and other substances we eat. Carbon dioxide is a waste product of animal life.

On the other hand, carbon dioxide is an essential food of plants. Through their leaves, especially, they absorb carbon dioxide from the atmosphere, decompose it. reject part of the oxygen, and store up the carbon in the form of complex compounds, such as starch. The sunlight and the green coloring matter (called chlorophyll) aid the plant in the formation of these com- pounds.

The relation of carbon dioxide to life is clear. Plants absorb carbon dioxide and transform, it into starch, whereas animals eat starch as food, assimilate it. and oxidize the carbon to carbon exhaled into the atmosphere ready for and so on.

Fig. 16. — Experiment showing the absorp- tion of carbon diox- ide and liberation of oxygen by plants

dioxide, which is the plants again.

The fact that plants take up carbon dioxide and reject oxygen can be readily illustrated, as shown in Fig. 16. Fresh green leaves are put into the flask, which is then completely tilled with water satu- rated with carbon dioxide. The stopper with its funnel is pushed in to exclude the air. the funnel is partly filled with the same liquid, and the test tube is filled and arranged as shown in the figure. On ex- posure to the sunHght for several hours, a gas collects in the test tube. The usual test shows that the gas is oxygen (21). Xo oxygen is produced if water free from carbon dioxide is used.

PRACTICAL CHEMISTRY

The significant relation of carbon dioxide and oxygen to plants and animals, which is often spoken of as the cycle of carbon and oxygen, is shown in Fig. 17.

Carbon dioxide in the air

Animals

Carbon dioxide in the air

Fig. 17. — Cycle of carbon (.4) and oxygen (B)

40. Carbon dioxide and fire extinguishers. — Carbon dioxide does not burn, but extinguishes burning substances. Instead of the gas itself, a saturated solution is frequently used to put out small fires. The solution is prepared, as

r~i ^m^

r r

—JMt^§mSSSS

'^^'^- ■

Fig. 18. — Special type of motor chemical engine for extinguishing forest fires and grass fires. Two large generating tanks are on top and portable extinguishers are on the running board

needed, in portable fire extinguishers and in chemical en- gines (Fig. 18) by the interaction of sulphuric acid and sodium bicarbonate. The ordinary fire extinguisher con- tains a solution of sodium bicarbonate and a loosely stop-

CARBON AND ITS OXIDES

percd bottle of sulphuric acid ; upon inverting the tank, the stopper of the acid bottle falls out, the two liquids mix, and the pressure of the generated gas forces the satu- rated solution of carbon dioxide out of the nozzle of the extinguisher ; some of the carbon dioxide itself escapes (Fig. '19). The water solu- tion of carbon dioxide together with the gas forces the oxygen of the air away from the fire and thereby reduces or en- tirely prevents combustion.

41. Carbon monoxide differs from car- bon dioxide. — Carbon monoxide is a compound of carbon and oxygen, but the two compounds differ in properties and composition. They can be prepared from each other by chemical processes.

Carbon monoxide, Hke carbon dioxide, is a gas without color, odor, or taste. But in other properties the two gases differ. Thus, carbon monoxide is only slightly soluble in water, and does not form a compound with water. It is lighter than carbon dioxide; a hter weighs 1.25 gm.

42. Carbon monoxide is poisonous. — Carbon monoxide is a dangerous poison because the lack of odor prevents its detection. A small fraction of a per cent of this gas in the air produces a stupefying effect. Many deaths have been caused by breathing air containing it. Carbon monox- ide impoverishes the blood by forming a compound with one of its constituents, and persons who have been poisoned by this gas cannot usually be revived by air, as in the case of suffocation by carbon dioxide. It is an ingredient of ordinary illuminating gas, and care should always be taken

Fig. 19. — Portable fire extinguisher (partly open) showing stop- pered acid bottle in original posi- tion

34 PRACTICAL CHEMISTRY

to prevent the escape of illuminating gas (as well as the gas from a coal stove or furnace) into rooms occupied by human beings. The exhaust gases from an automobile engine contain carbon monoxide, and special care should be taken to ventilate a small garage. The pulmotor is often used to revive persons who have been overcome by gases containing carbon monoxide (29).

43. Chemical conduct of carbon monoxide. — Carbon monoxide, unhke the dioxide, burns in air or oxygen. The flame is blue. The product is carbon dioxide. The equa- tion for this chemical change is : —

Carbon Monoxide + Oxygen = Carbon Dioxide

(Carbon-Oxygen) (Carbon-Oxygen)

The flickering bluish flame often seen on the top of a coal fire is caused by the burning carbon monoxide.

Not only does carbon monoxide unite readily with oxygen, but it withdraws oxygen from hot oxides ; carbon itseK acts in the same way. In chemical language carbon monoxide (and also carbon) is a reducing agent; i.e. it withdraws oxygen from compounds. This chemical re- moval of oxygen is called reduction. In the manufac- ture of iron from iron ores, the ore, which is an oxide (FeoOs), is reduced by carbon monoxide in a blast furnace, the gas for this purpose being produced by the incomplete combustion of coke. The same kind of chemical change takes place easily with copper oxide and hydrogen (55).

44. The two carbon oxides are formed in a coal fire. — The oxygen of the air entering at the bottom of a coal fire combines with the hot carbon of the coal and forms carbon dioxide — the first change. But the carbon diox- ide in passing up through the upper layer of hot coal is reduced by the carbon to carbon monoxide — the second

CARBON AND ITS OXIDES

20. — Carbon oxides are formed in a coal fire

change. The carbon monoxide escapes through the top of the tire into the air, where much, or all, of it burns to carbon dioxide — the third change. Therefore we see that the two oxides are closely re- lated chemically and pass into one another readily, especially in a coal fire (Fig. 20).

45. Formation and preparation of carbon monoxide. — Carbon mon- oxide is always formed when carbon and certain carbon compounds {e.g. gasolene) burn in a limited supply

of air. It is also formed when steam is passed through a hot fire of hard coal or coke. The gaseous product is a mixture of carbon mon- oxide and hydrogen ; this mixture if enriched by vapor from petro- leum oil so that it burns with a yellow flame, is called water gas and is used, alone or with other gases, as illuminating gas. Recall that the carbon monoxide makes such a gas poisonous.

If steam and air are together passed through hot carbon, the gaseous product contains nitrogen and some carbon dioxide besides carbon monoxide and hydrogen; it is called fuel or producer gas. It is easily made and liberates considerable heat in burning, and is there- fore used extensively as a fuel in industrial processes, e.g. in mak- ing open-hearth steel.

Carbon monoxide is usually prepared in the laboratory by heating a mixture of oxalic acid, or formic acid, and sulphuric acid, and col- lecting the gaseous product over water (294) .

46. Carbon dioxide and monoxide are good examples of compounds. — In 9 we learned that compounds have three essential characteristics. These are briefly: (i) The elements in compounds are united chemically, (2) the properties of compounds differ from the elements in them, and, most important of all, (3) compounds have a definite composition.

36 PRACTICAL CHEMISTRY

The chemical conduct, formation, and preparation of both oxides show how well these two compounds exhibit the first and second characteristics. Both require specific chemical action — usually involving heat as an agent for their formation, preparation, and chemical conduct. And obviously both differ in properties from carbon and oxygen, the elements that compose them.

The third characteristic, viz. constant composition, is not apparent from any experiment thus far described. The composition of a compound is the proportion by weight of each element in it. This proportion can be found by simple experiments, though the work must be done accurately. One method is to weigh the carbon and oxygen that com- bine to form carbon dioxide and also carbon monoxide. Another method is to decompose a known weight of each compound and weigh the carbon and oxygen obtained. We know from many accurate experiments that the propor- tion of carbon to oxygen in carbon dioxide is i to 2.66, and in carbon monoxide it is i to 1.33. These results mean, to repeat, that these two substances are composed of the elements carbon and oxygen chemically combined in a definite, unvarying proportion by weight. That is, these substances are chemical compounds.

Let us consider these figures more carefully. First, in the case of each compound the proportion is definite and unvarying. If it were not, these substances would not be compounds but mixtures ; in a mixture, you will re- member, the proportion of the ingredients may vary (7). Second, if we compare the two proportions, we notice they are different. If they were not, we should not have the two different compounds. Third, if we compare the num- bers expressing the oxygen (2.66 and 1.33), we notice a simple relation, i.e. one is exactly twice the other. This

CARBON AND ITS OXIDES 37

means, in other words, that the weight of the oxygen com- bined with the carbon in carbon dioxide is twice the weight combined with the carbon in carbon monoxide.

47. Law of multiple proportions. — There are many other cases hke carbon dioxide and carbon monoxide. That is, there are groups of compounds of the same elements, and in any one group there is a simple multiple relation (1:2, 2:3, 1:3, etc.) between the weights of the element that are combined with a fixed weight of the other element. (This is fully worked out in 99.) The general fact of mul- tiple relations is sometimes summarized in a brief form called the law of multiple proportions, which may be stated thus : —

In a group of compounds of the same elements, small whole numbers will express the different weights of one element that combine with a fixed weight of the other clement.

This law and the law of definite composition (9 and this section above) are important laws in chemistry. Later we shall study them more fully. (See especially Chapter VII.)

EXERCISES

1. Prepare a brief summary of this Chapter (a) in the form of short topics, and {h) as a connected narrative.

2. Prepare a summary of Chapters I, II, III showing ((/) the re- lations of elements and compounds, (/>) the properties of compounds, and (f) the laws (4, 9, 47).

3. In what form does free carbon occur in nature ? Name ten familiar solids, three liquids, and two gases which contain carbon.

4. Name several ways in which carbon dio.xide is formed.

5. How is carbon dioxide prepared in the laboratory? State the properties of carbon dioxide. Describe its chemical conduct.

6. How is carbon dioxide related to plant and animal life?

7. Describe a fire extinguisher.

8. Compare carbon dioxide and carbon monoxide.

9. How are the two carbon oxides related to a coal fire?

SS PRACTICAL CHEMISTRY

10. Describe fully the action of carbon dioxide on calcium hydroxide. State the reaction by an equation.

11. What is the test for (a) carbon, (b) carbon monoxide, (c) car- bon dioxide, (d) a carbonate?

12. State the equation for (a) the oxidation of carbon to carbon mon- oxide and (b) the reduction of carbon dioxide to carbon monoxide.

13. Illuminating gas, water gas, and the gas that escapes from a coal fire are poisonous. Why? What is a pulmotor and for what is it used?

14. State and illustrate the law of (c) constant (or definite) com- position and (b) multiple proportions.

15. Define and illustrate reduction. Name two reducing agents studied in this chapter.

16. Make a list of the new chemical words in this chapter and de- fine each.

17. Make a list of the compounds mentioned in this chapter and state as fully as possible the elements in each.

PROBLEMS

1. What is the weight of lo liters of carbon dioxide gas? Of io,ooo cc? Of3l.?

2. A pupil prepared 5 bottles of carbon dioxide gas each holding 250 cc. How many gm. of carbon dioxide were prepared?

3. Marble, if treated with acid, yields about 44 per cent of its weight as carbon dioxide. If 0.5 kg. of marble was used, how many (a) gm. and (b) 1. of carbon dioxide were formed ?

4. What weight of marble corresponds to 150 1. of carbon dioxide?

CHAPTER IV HYDROGEN

48. Occurrence of hydrogen. — Hydrogen, like oxygen, is a gaseous element. But it does not occur free to any great extent in nature ; natural gas — a mixture of com- bustible gases which issues from the earth in certain locaU- ties — contains about 2 per cent. Mixtures of gases for Hghting and heating, e.g. water gas and producer gas (45), contain from 35 to 50 per cent of hydrogen. Combined hydrogen is a constituent of many important compounds. Water is 11. 18 per cent of hydrogen. The human body is about 10 per cent of hydrogen. Hydrogen is a constituent of all acids and bases — important classes of compounds to be studied later. (See Chapter XII.)

All plants and animals contain compounds of hydro- gen with carbon and oxygen, and in some cases with nitro- gen also.

Compounds of hydrogen and carbon form a large and im- portant class of compounds called hydrocarbons, which are ingredients of petroleum (and its products, such as kero- sene, gasolene, paraffin, and lubricating oils), illuminating gas, and producer gas.

49. Hydrogen is prepared from acids or water. — Hydro- gen is readily and conveniently prepared by the interaction of certain metals and acids. The metals are usually zinc, iron, or magnesium, and the acids are dilute water solu- tions of sulphuric acid (H2SO4) or hydrochloric acid (HCl). The hydrogen is liberated from the acid, and the metal com-

PRACTICAL CHEMISTRY

bines with the rest of the acid to form a compound which usually remains dissolved in the liquid. On a large scale a Kipp apparatus (Fig. 21) is sometimes used.

In the laboratory hydrogen is usu- ally prepared in a small generator, and collected over water in a pneumatic trough (Fig. 22). Zinc is put in the bottle A and acid is introduced through the dropping tube B by pressing the clamp. The hydrogen passes out through the delivery tube D into the pneumatic trough, bubbles up into the bottles, and displaces the water. No flame should be near during the preparation of hydrogen, because mix- tures of air and hydrogen explode vio- lently when ignited.

Hydrogen can also be prepared by allowing certain metals and water to interact. Sodium interacts less rapidly than potassium. Calcium interacts slowly with water. But potassium

Fig. 21. — Kipp ap- paratus for generat- ing hydrogen

interacts so rapidly that the heat ignites the liber- ated hydrogen (Fig. 23). If a small piece of sodium is dropped upon cold water, the sodium melts into a shining globule, which spins about rapidly on the water with a hissing sound and finally disappears with a slight explosion. If the

Fig. 22, — Apparatus for preparing hydrogen in the laboratory

HYDROGEN

Fig. 23. — The interac- tion of water and potas- sium produces enough heat to ignite the liber- ated hydrogen

sodium is wrapped in a piece of fine wire gauze, or

of tea lead pierced with small holes, and dropped into a

dish of water, the hydrogen gas can be collected in an

inverted test tube full of water (Fig. 24). Hydrogen, together with oxygen,

is liberated from water by passing

a current of electricity through

water containing sulphuric acid or

sodium, hydroxide (16, 88).

Hydrogen can also be prepared

by passing steam — the gaseous

form of water — over heated metals

(Fig. 25). This experiment was

first performed by Lavoisier, in

1 783, while he was studying the com- position of water. He passed steam

through a red-hot gun barrel containing bits of iron.

The oxygen of the steam combined with the iron, and the hydrogen escaped from the tube. Since Lavoisier was then studying the composi- tion of water and not especially the prop- erties of hydrogen, he naturally thought of the gas as essential for forming water. So he named the gas hydrogen, which means literally " water former."

50. Hydrogen can be prepared from alka- lies. — Hydrogen can also be prepared by boiling solutions of certain alkalies with some elements. Thus, if sodium hydroxide (NaOH) is boiled with aluminium or silicon, hydrogen is liberated.

51. A new kind of chemical change illustrated by the preparation of hydro- gen. — The preparation of hydrogen by

Fig. 24. — Appa- ratus for collect- ing hydrogen liberated by the interaction of water and so- dium

PRACTICAL CHEMISTRY

the interaction of a metal and an acid, water, or an alkali illustrates a third kind of chemical change, viz. substitu- tion, or, as it is sometimes called, displacement or replace-

T^^^

Fig. 25. — Modern form of Lavoisier's apparatus for showing the formation of hydrogen by the interaction of steam and heated iron

ment. In the case of zinc and sulphuric acid, zinc is sub- stituted chemically for hydrogen, i.e. the hydrogen is dis- placed from the acid by the zinc. This chemical change can be expressed by the following equation : —

Zinc -\- Sulphuric /Vcid = Hydrogen + Zinc Sulphate

(Hydrogen-Sulphur- (Zinc-Sulphur-

Oxygen) Oxygen)

Similarly we have : —

Sodium 4- Water = Hydrogen -|- Sodium Hydroxide

(Hydro- (Sodium-Hydrogen-

gen-Oxygen) Oxygen)

Iron + Water = Hydrogen -\- Iron Oxide

(Hydro- (Iron-Oxygen)

gen-Oxygen)

Aluminium + Sodium Hydroxide = Hydrogen + Sodium Aluminate

(Sodium-Hydrogen- (Sodium- Aluminium-

Oxygen) Oxygen)

HYDROGEN 43

We define substitution as a chemical change in which one element displaces another in a compound.

52. Some properties of hydrogen. — Hydrogen has no taste or color. The pure gas has no odor, though hydrogen as ordinarily prepared has a disagreeable odor, due mainly to impurities in the metals used. Hydrogen is very slightly soluble in water, less so than oxygen.

Hydrogen is the Ughtest known substance. Volume for volume hydrogen is about one fourteenth as heavy as air and one sixteenth as heavy as oxygen. One hter at 0° C. and 760 mm. weighs only 0.09 gm. (exactly 0.0898 gm.). Hydrogen, being so light, diffuses rapidly, i.e. it quickly passes through porous substances (e.g. unglazed porcelain, rubber, and thin sheets of metal), mixes rapidly with other gases, and freely escapes into space in all directions.

The extreme lightness as well as the rapid diffusion of hydrogen can be shown by simple experiments. If a bottle of hydrogen is ex- posed to the air a minute or two, and a lighted match then dropped in, the match merely burns ; if hydrogen were present, a loud explo- sion would have occurred. Or, if a bottle of hydrogen is held just beneath a bottle of air, the gases quickly change places more or less completely, the hydrogen rising into the upper bottle and forcing out some of the air ; if the experiment is well done, the gas in the upper bottle burns or explodes gently, but usually the loud explosion shows that only a part of the hydrogen flowed into the upper bottle.

53. Chemical conduct of hydrogen. — At ordinary tem- peratures and under ordinary conditions, hydrogen, like oxygen, is not an active element. A mixture of hydrogen and oxygen can be kept indefinitely at the temperature of the laboratory. But if the mixture is heated to about 800° C. or if a flame is brought very near it, the gases unite with a violent explosion.

Under special conditions hydrogen unites with oxygen quietly, i.e. hydrogen can be made to burn quietly in oxygen

PRACTICAL CHEMISTRY

or air. Hydrogen is generated in the apparatus shown in Fig. 26 ; the gas passes through the drying tube and escapes in a fine stream through the small opening in the platinum tip. After all the air has been driven out of the whole

apparatus by the hydro- gen, the gas is Hghted by holding a lighted match at the end of the tip. The hydrogen burns in the air with an almost invisible but very hot flame. If a small, dry, cold bottle is held over the flame, 26. — Apparatus for burning water vapor is deposited hydrogen inside the bottle. Water

is the product of the combustion of hydrogen. That is, hydrogen in burning unites with oxygen, just as other burning substances do. Thus :

Hydrogen -f Oxygen = Water This chemical change is an example of combination and also of oxidation (22). The two elements, hydrogen and oxygen, unite to form the compound water ; hydrogen is oxidized, and the product might be called hydrogen oxide, though it is called by its more familiar name water.

The film of water that may be seen on the bottom of a vessel placed over a lighted gas range or a Bunsen burner is the condensed vapor formed by the burning hydrogen, and hydrogen compounds, of the illuminating gas. Organic substances containing hydrogen, such as wood and paper, when burned, yield water as one of their products.

Although a small jet of hydrogen burns quietly in air or in oxygen, a mixture of hydrogen and air burns so rapidly that the combustion is practically an explosion. There- fore, the air should be fully expelled from the apparatus

HYDROGEN 45

in which hydrogen is being generated and all leaky joints should be tightened before the gas is collected ; no flames, large or small, should be near. Neglect of these precau- tions has caused serious accidents.

The chemical conduct of hydrogen with the element chlorine is similar to that with oxygen. If the two gases are mixed in the dark, they do not react. But if a mixture of hydrogen and chlorine is heated or exposed to the sunlight, the two gases combine with ex- plosive violence. However, a hydrogen flame, if lowered into a bottle of chlorine, continues to burn quietly. In this case, as with oxygen, the two elements, hydrogen and chlorine, unite ; the product is the compound hydrogen chloride. The burning of hydrogen in chlorine illustrates an extension of the term combustion. No oxygen is involved, but it is a case of chemical combination accompanied by light and heat. (See 25.)

Certain metals under special conditions hasten the rate at which hydrogen combines with elements. Thus, if a mixture of hydrogen and oxygen is passed over finely di- vided platinum, the gases, which ordinarily would scarcely combine at all, now begin to react to form water. Similarly, hydrogen and nitrogen can be made to unite into the com- pound ammonia, if the mixture is passed under proper conditions over certain metals. The platinum and the other metals, as far as we know now, do not undergo a chemical change in these reactions. They hasten a very slow chemical reaction. A metal w^hich acts thus is called a catalyst or a catalytic agent. Its function is sometimes compared to that of lubricating oil on a machine.

54. Hydrogen does not support combustion. — Hydro- gen burns, but does not support combustion. These facts are illustrated by putting a hghted taper into an inverted bottle of hydrogen (Fig. 27). The taper ignites the hydrogen, which burns at the mouth of the bottle. The taper does not burn inside the bottle, whereas when it

0 PRACTICAL CHEMISTRY

is slowly withdrawn through the burning hydrogen it is relighted.

55. Hydrogen is a reducing agent. — Hydrogen not only combines energetically with free oxygen, but it also

^ ^ withdraws oxygen from compounds. This

chemical removal of oxygen is called re- duction, and the substances that remove the oxygen are called reducing agents. (Compare 43.) Hydrogen is a vigorous reducing agent, just as oxygen is an ener- getic oxidizing agent. When oxides of ^T , certain metals are heated in a current

Fig. 27. — Hydro- gen burns but of hydrogen, the oxygen of the oxide is

does not support chemically removed and combines with the combustion hydrogen^ to form water; the metal is

left uncombined. Thus, by heating copper oxide in hydro- gen, water and metalHc copper are produced. Chemically speaking, the copper oxide is reduced by the hydrogen. The chemical change is substitution (the hydrogen being substituted chemically for the metal), and it can be ex- pressed thus : —

Copper Oxide + Hydrogen = Water + Copper

(Copper-Oxygen) (Hydrogen-Oxygen)

This chemical change can be interpreted from the standpoint of oxidation, because the hydrogen is oxidized to water at the same time the copper oxide is reduced. In fact, the processes of reduction and oxidation are closely related and either one may be emphasized in interpreting the chemical change. In its simplest form, reduction is the opposite of oxidation. Later we shall see that the terms oxidation and reduction are both used in a broader sense.

56. Test for hydrogen. — A simple test for hydrogen is that it extinguishes a small flame, such as a blazing taper or joss stick, but

HYDROGEN 47

is lighted at the same time, often with an explosion, and continues to burn until the gas is exhausted. A conclusive test is that it burns with a hot flame and forms water as the sole product.

Fig. 28. — Welding an iron grill with an oxy-acetylene flame

57. Uses of hydrogen. — On account of its extreme light- ness, hydrogen is used to fill balloons and dirigible airships ; small balloons are usually filled with coal gas and larger dirigible craft will, it is predicted, be filled with the in- combustible gas helium (127).

The intense heat of the hydrogen flame is utilized in the oxy-hydrogen blowi^ipe. The essential part of the burner is two pointed metal tubes. The inner and smaller one is for oxygen, and the outer and larger one for hydrogen; the gases are forced out of these small openings by the pres- sure maintained in the storage tanks. The temperature of the oxy-hydrogen flame is about 2000° C. The flame is used to melt platinum. When the flame strikes against a piece of lime, the latter becomes intensely bright. Thus used, it is called the hme or calcium Hght, and is utihzed in the stereopticon. The oxy-hydrogen flame has been largely replaced by the electric furnace and the oxy-acety-

PRACTICAL CHEMISTRY

Fig. 29. — A blast lamp

lene flame (Fig. 28) as sources of intense heat. (Acetylene is a compound of carbon and hydrogen — C2H2.)

A substitute for the oxy-hydrogen flame is used in

the laboratory. It is called a blast lamp (Fig. 29). Illumi- nating gas (which contains hydrogen and hydrocarbons) and air are used in- stead of hydrogen and oxygen.

The most extensive use of hydrogen is in transforming certain oils, e.g. cotton-seed oil, into sohd edible fats. This process is called hydrogenation (363).

The hydrogen needed for these uses is prepared by the electrolysis of water (16).

EXERCISES

1. Prepare a summary of this chapter.

2. Compare oxygen and hydrogen in tabular form.

3. How can hydrogen be distinguished from (a) oxygen, (b) car- bon monoxide, (c) carbon dioxide, (d) air?

4. Summarize the conspicuous properties of hydrogen. Describe its chemical conduct.

5. Why is there danger of an explosion in generating hydrogen? How can the danger be avoided?

6. Define and illustrate (a) reduction and (b) reducing agent. Com- pare with (a) oxidation and {b) oxidizing agent.

7. Make a list of the new chemical words in this chapter and de- fine each.

8. In using hydrogen for balloons, what property of the gas might cause disaster?

9. Topics for home study, (a) Burning of hydrogen, (b) Re- lation of hydrogen to water, (c) Lavoisier's experiment, [d) Re- view of chemical change, (e) Uses of hydrogen.

HYDROGEN 49

PROBLEMS

1. Sulphuric acid contains 2.04 per cent of hydrogen, (a) How many grams must be decomposed to yield 85 gm. of hydrogen (at 0° C. and 760 mm.)? (b) 85 cc?

2. Hydrochloric acid contains 2.74 per cent of hydrogen. How many tons must be decomposed to yield a ton of hydrogen?

3. Water contains 11. 18 per cent of hydrogen. How many grams of hydrogen can "be prepared from 230 gm. of water? How many liters (at 0° C. and 760 mm.) ?

4. (a) How many liters (at 0° C. and 760 mm.) will 45 gm. of hydro- gen occupy? (b) How many grams will 45 1. weigh?

• 5. A student prepared enough hydrogen (at 0° C. and 760 mm.) to fill six bottles, each holding 250 cc. (a) How many grams were prepared? (b) How much sulphuric acid was used up?

6. A cylindrical tank i m. long and 25 cm. in diameter is filled with hydrogen (at 0° C. and 760 mm.). How many grams does the gas weigh ?

7. A hydrogen manufactory produces daily 19,000 cubic meters of gas. (a) Express this volume in liters and in cc. (b) What would this volume weigh, if measured at 0° C. and 760 mm.?

8. A Zeppelin balloon had a capacity of 351,150 cubic feet. What weight of hydrogen was needed to fill it? (Assume i cu. m. = 35.32 cu. ft. ; also that the volume of hydrogen was measured at 0° C. and 760 mm.)

CHAPTER V THE MEASUREMENT OF GASES

58. Introduction. — So far we have studied four gases : the two elements, oxygen and hydrogen, and the two com- pounds, carbon dioxide and carbon monoxide. We shall study many other gases in the succeeding chapters.

59. What is our problem? — In chemistry we often need to know the weight of a given volume of gas, or to compare the weights of the same volume. But we cannot deter- mine the weight of a gas nor compare weights until we agree on the conditions under which we are to measure the vol- ume. Thus, when we say one hter of oxygen wxighs 1.43 gm., we really mean that the hter has this weight if the vol- ume is measured at a certain temperature and pressure (0° C. and j6o mm.).

All gases change in volume with changes in tempera- ture and pressure. In this chapter we shall study: (i) the effect of changes of tem.perature and pressure on the volume of a gas, (2) how the required volume of a gas is found at the standard temperature and pressure, and (3) how the weight of a given volume of a gas is found.

60. Change in the volume of a gas with changes in tem- perature. — Gases expand when heated and contract when cooled. This change in volume is uniform for all gases, pro- vided of course there is no change in pressure. Thus, if the pressure is kept constant, 273 cc. of a gas at 0° C. be- comes 274 cc. at 1° C, 280 cc. at 7° C, 272 cc. at -1° C, 266 cc. at -7° C, and so on. That is, for each degree of

THE MEASUREMENT OF GASES

CENTIGRADE AliSOLUTE Water

100^

O''

•182.5^

Boils

Water

Freezes

Oxytren

Boils

Absolute

3730

273°

change in temperature the volume at o"^ C. changes ar^- This general fact is stated as the law of Charles, thus : —

Gases under constant pressure change in volume uniformly with the same change in temperature.

At this rate of contraction a gas would have no vol- ume at —273° C.I But this rather startUng result does not actually happen, because gases become liquids above this temperature. Nevertheless this point(— 273°) on the centigrade thermometer is useful, for it serves as the starting point of a temperature scale called absolute tem- perature. The temperature — 273° C. is called absolute zero, and tempera- tures reckoned from this point as zero are called absolute temperatures. If an "absolute thermometer" were constructed, and compared wdth a cen- tigrade thermometer, the scales would be seen to be simply related (Fig. 30). If we designate absolute temperatures by A. and centigrade by C, then we see that 273° A. is 0° C, 373° A. is 100° C, and 0° A. is —273° C. That is, absolute degrees are obtained by adding 273 to centigrade degrees.

How is absolute temperature used ? Suppose, as above, we have 273 cc. of gas at 0° C. This volume becomes 274 cc. at 1° C. and 280 cc. at 7° C. Now 0° C, 1° C, and f C. are 273° A., 274° A., and 280° A. That is, stated in general terms, the volumes of a given weight of gas at different temperatures are in the same ratio as the absolute temperatures.

90.5°

Ze;

0°

Fig. 30. — Centigrade and absolute scales

52 PRACTICAL CHEMISTRY

61. How we apply the law of Charles. — Suppose we have lOO cc. of oxygen at 25° C. and we wish to know what the volume would be at 0° C. It is not necessary to cool the gas to 0° C. and measure the volume. We know : (i) that a gas contracts uniformly for each degree of de- crease in temperature, and (2) that the volumes are in the same ratio as the absolute temperatures. Hence we can compute the volume to which the 100 cc. would contract if it were cooled from 25° to 0° C. The steps in the pro- cess are these : {a) change the centigrade temperatures to absolute temperatures by adding 273 to each; (b) make 273 (i.e. o -f 273) the numerator of a fraction and the other sum the denominator ; (c) multiply the volume by this fraction. Thus: (a) o + 273 = 273, and 25 + 273 = 298; (b) Hi; (c) 100 X Ui = 91.61. That is, 91.61 cc. is the volume 100 cc. would become if cooled from 25° C. to o^ C.

In solving problems involving change in temperature, care should be taken to notice the conditions. For example, 273 is the numerator only when the final temperature is 0° C. ; if the final temperature is another number, the numerator is this number + 273. In checking results, it should be remembered that decrease in temperature means decrease in volume.

62. Change in the volume of a gas with changes in pres- sure. — Many gases are collected in bottles or tubes over water or mercury in a pneumatic trough or tall jar. In measuring the volume of the gas, the levels of the liquid are made the same inside and outside the bottle or tube by raising or lowering it. The gas thus becomes under at- mospheric pressure, i.e. the pressure of the atmosphere exerted on the exposed surface of the liquid in the trough or jar is transmitted through the water to the gas. Hence the pressure which the gas is under is the same as the pres- sure of the atmosphere when the gas volume is read.

THE MEASUREMENT OF GASES

■SOO

7.-;^)

700

The pressure of the atmosphere is found merely by read- ing the barometer. A common form of barometer is shown in Fig. 31. It consists of a small-bore, strong glass tube, about one meter long, and closed at the upper end. The tube is nearly full of mercury and the open, lower end dips into a small reservoir of mercur>\ The space above the mercury at the upper end of the tube is a vacuum. As the pressure of the atmosphere changes, the column rises or falls. Near the top of the supporting board back of the tube is a scale by which we can tell the height of the mercury column above the surface of the mer- cury in the reservoir. That is, by reading on the scale the height of the mercury; — '' reading the barometer " — we are measuring the pres- sure of the atmosphere at the time of the ob- servation. The normal, or average, height of the barometer is 760 millimeters (mm.).

As stated in 59 the volume of a gas changes with change in pressure. The volume con- tracts if the pressure is increased and expands if the pressure is decreased. The effect of pres- sure on the volume of a gas was first studied by Boyle (Fig. 32). Fie found that 100 cc. of air, for example, contracts to 50 cc. if the pressure is doubled, and also that 100 cc. expands to 200 cc. if the pressure is halved. This relation of volume to pressure, which applies to all gases, is sometimes stated as the law of Boyle, thus : —

The volume of a gas at constant temperature changes in- versely with changes in pressure.

63. How we apply Boyle's law. — In most of the experi-

Fig. 3I--A barometer

PRACTICAL CHEMISTRY

ments in chemistry we actually read a gas volume at the prevailing pressure, although we need to know what the volume would be if the pressure were 760 mm. As in the

case of temperature and volume, it is not neces- sary (and certainly not convenient !) to wait until the barometer be- comes 760 mm. We know that volumes are inversely related to their corresponding pressures (62) . Hence we can com- pute the volume which 100 cc. of gas, for ex- ample, would have if the pressure were changed from 775 mm. to 760 mm. The steps in the process are these : {a) Make 760 the denominator of a frac- tion and 775 its numerator; (b) multiply the volume by this fraction. Thus: (a) -fjf ; (b) 100 X fe-l" = 101.97.

Fig. 32. — Boyle (1626-1691

In solving problems involving change in pressure, care should be taken to notice the conditions. (Compare end of 61.) For example, 760 is the denominator of the fraction only when the final pressure is the normal pressure (760 mm.) ; if the final (or desired) pressure is another number, the denominator is this number. It is advisable also to check the result by comparison with the original volume, remembering that decrease in pressure means increase in volume.

64. How we reduce a gas volume to standard conditions.

— Measuring gases usually involves reading both tempera- ture and pressure. The normal or standard temperature selected in chemistry for measuring gases is 0° C, and the

THE MEASUREMENT OF GASES

normal or standard pressure is 760 mm. These conditions (o° C. and 760 mm.) are called normal or standard con- ditions. It is inconvenient, often impossible, to measure gas volumes at exactly 0° C. and 760 mm. So it is cus- tomary to measure the volume at the prevailing tempera- ture and pressure, and then compute the volume the gas would have if it were at 0° C. and 760 mm. This mathe- matical calculation is called reducing to standard conditions or correcting the volume for temperature and pressure. The final volume is called the corrected volume. Let us take an example. Suppose a gas volume measures 100 cc. at 25° C. and 775 mm., and we wish to reduce this observed volume to standard conditions. We apply the laws of Charles and Boyle, and merely combine the two mathe- matical processes already described (61, 63). Thus : ■—

lOO A T98" X Te^o^

9341

Fig. 7,7,. — Apparatus for finding the weight of a liter of oxygen

65. How we find the weight of a given volume of a gas. — In preceding sections the weight of one liter of different gases is given. Thus, oxygen is 1.43, carbon dioxide is 1.98, carbon monoxide is 1.25, and hydrogen is 0.09 gm.

56 PRACTICAL CHEMISTRY

Let us see how the weight of one Uter of oxygen is found, (a) The apparatus is shown in Fig. 33. A mix- ture of potassium chlorate and manganese dioxide is put in the test tube A. The part AF is weighed. The bottle B is filled with water. The empty bottle D is weighed. The mixture is heated and oxygen forces water from B into D. When sufficient gas has been liberated, the heat- ing is stopped and the apparatus is allowed to cool. (b) The temperature and pressure are read — say 20° C. and 755 mm. The levels in B and D are made the same (by raising one bottle). The water in D is measured ; its vol- ume is the same as the volume of the oxygen — say 1.75 liters. AF is weighed ; its loss is the weight of the oxygen — say 2.322 gm. (c) The observed volume is reduced to stand- ard conditions; thus: 1.75 X ffl X fro = 1-625. (d) The weight of one Uter of oxygen at 0° C. and 760 mm. is found by dividing the weight of the oxygen by its corrected volume; thus, 2.322 -7- 1.625 = 1.429.

We shall see later that a slight modification must be made in the value of the pressure when the oxygen is collected over water (74).

66. The nature of gases. — We have seen in this chapter that gases change uniformly in volume with changes in temperature and pressure. They behave alike in many other ways, e.g. they diffuse readily. The similarity in properties suggested long ago that gases are fundamentally alike. We beheve that gases are made up of minute parti- cles, too small to be seen with a microscope. These parti- cles are called molecules. The molecules of a gas are far apart, i.e. the distances between molecules are large com- pared with the molecules. They are exceedingly small particles in a large space. Furthermore, we believe that the molecules are in constant motion, i.e. flying about

THE IMEASUREMENT OF GASES 57

rapidly in all directions, striking each other and the walls of the containing vessel, rebounding, and so on ceaselessly. Moreover, the movements, so it is believed, increase in velocity as temperature rises, and decrease as it falls.

These suppositions are part of the kinetic-molecular theory of gases. We cannot see the moving (kinetic) mole- cules, but many facts lead us to beheve in their existence. We state our belief as a theory, i.e. a supposition or mental picture which enables us to interpret related facts.

The law of Boyle seems much clearer if we interpret it by the kinetic-molecular theory. Thus, when a gas is compressed, the molecules are pressed more closely to- gether. And when a gas is compressed to half its volume, the pressure produced by the blows of the moving mole- cules is doubled because the number of blows per second against the walls is doubled.

Similarly, the law of Charles is more readily understood if we think of the m.olecules of a gas as moving more rapidly, and consequently requiring more space when the temperature is raised.

Again, the constant m^otion of the molecules helps us understand the rapid movement of gases called diffusion, e.g. the rising of illuminating gas throughout a building from a small leak in the cellar.

More will be said later about the nature of gases ; the nature of liquids, too, will be considered. (See 86 and 101.)

EXERCISES

1. Prepare a summary of this chapter.

2. (a) State Boyle's law. Illustrate it. (b) State Charles's law. Illustrate it.

3. Give examples from everyday life of {a) expansion and of (b) con- traction of gases caused by change of temperature.

4. Apply E.xercise 3 to change of pressure.

58 PRACTICAL CHEMISTRY

5. (c) Change these centigrade readings to absolute : loo, o, IS) ~i5) 250, 273, —273. (b) Change these absolute readings to centigrade: 273, o, 200, 100, 473, 180, 373.

6. (a) When a given volume of gas is reduced to standard con- ditions, is the weight of the gas changed? (b) Is the gas itself actually reduced to 0° C. and 760 mm.?

7. From the data given in 65, calculate the weight of a liter of oxygen without making the correction for temperature and pressure. Compare the result with the correct weight.

8. Topics for home study, (a) Boyle's contributions to science. (b) Thermometers, (c) Barometers, (d) Atmospheric pressure.

PROBLEMS

1. Reduce the following to the volume occupied at 760 mm. : (a) 20 cc. at 745 mm. ; (b) 45 cc. at 765 mm. ; (c) 450 cc. at 755 mm. ; (d) 1.5 1. at 763 mm.; (e) 2.5 1. at 745 mm.; (/) 500 cc. at 75 cm.; (g) 76 cc. at 76 cm. ; (h) 900 1. at 749 mm.

2. Reduce the following to the volume occupied at 0° C. : (a) 170 cc.at8o°C.; (b) 450 cc. at 15° C. ; (c) 70.6 cc. at 17° C. ; (d) 49 cc. at 19° C. ; (e) 356 cc. at 34° C. ; (/) 48 cc. at 27° C.

3. Reduce the following to the volume at standard conditions : (a) 250 cc. at 780 mm. and 20° C. ; (b) 140 cc. at 745 mm. and 21° C.

4. Reduce the following to the volume at standard conditions : (a) 247 cc. at 720 mm. and 14° C. ; (b) 1000 cc. at 750 mm. and 18° C. ; (c) 1480 cc. at 765 mm. and 81° C.

5. A volume of oxygen measured 375 cc. when the barometer was 740 mm., and the thermometer was 27° C. What would be the volume at the standard pressure and temperature? Ans. 332.3 cc.

6. Find the weight of 29 cc. of oxygen at 23° C. and 776 mm. Ans. 0.039 gm.

7. Calculate the weight of hydrogen in a vessel of 10 liters capacity, filled when the barometer reads 756 mm. and the thermometer 18° C.

8. If a volume of hydrogen measures 100 cc. at 14° C. and 755 mm., what will be its volume when measured at —8° C. and 780 mm.?

9. A gas measures 637 cc. at 755 mm. and 17° C. Find its volume at 730 mm. and 30° C.

10. Correct 250 cc. of oxygen at 18° C. and 745 mm. for tem- perature and pressure. Ans. 229.9 cc.

11. (a) What is the volume of 200 gm. of hydrogen at 0° C. and 760 mm.? (b) What would be its volume at 15° C. and 750 mm.?

CHAPTER VI WATER

67. Importance and occurrence of water. — Water is one of the most familiar of natural substances. Its varied properties and common uses make it an indispensable sub- stance. It is a compound of hydrogen and oxygen.

Water is always present in the air as a vapor, which is continually condensing into rain, clouds, mist, fog, dew, hail, frost, and snow. W^ater occurs in vast quantities on and beneath the surface of the earth as a liquid ; and also as a solid (snow and ice). The soil and plants contain con- siderable water. Many common foods, especially vege- tables, contain from 40 to 90 per cent of water. The human body is nearly 70 per cent water.

68. Natural waters are not pure. — Water is never found pure in nature. Even rain water, which is usually regarded as the purest natural water, contains dust and gases washed from the air. Water which flows along the surface of the earth or underground dissolves substances from the rocks and soil. Hence river water contains earthy impurities brought by the underground and surface water ; it is also often contaminated with compounds formed by the decom- position of animal and vegetable matter, the so-called ''or- ganic matter," or with refuse from manufactories.

Ocean water contains a large proportion of impurities, especially mineral substances Hke common salt and similar compounds of calcium and magnesium. The peculiar taste of ocean water is due chiefly to the presence of these substances.

69. Purification of water. — Since natural waters are not pure, it is often necessary to remove certain impurities to

PRACTICAL CHEMISTRY

make the water suitable for the desired use. Besides dis- solved mineral matter and organic matter, water may con- tain suspended matter such as fine particles of clay or other earthy substances and partly disintegrated organic sub- stances. Water, too, often contains sewage or organic matter and bacteria more or less associated with sewage. Water used for drinking and cooking should be as pure as possible, and it is desirable as well as economical that water be purified sufficiently for all practical purposes. The water of some towns and cities is purified by filtering it slowly on a large scale through layers of sand and gravel. Such a filter remioves suspended matter almost completely, though it must be frequently cleaned. In some locaUties the water is stored in a large settling basin or reservoir ; here the suspended soHd matter slowly settles, the process sometimes being hastened or aided by adding alum, which forms a sticky substance to which particles cling. Some- times the water is freed from organic matter by spraying it into the air (Fig. 34). This is an oxidation process.

Fig. 34. — Purification of water by spraying it into the air. An aerator of the New York City water system by which 376 million gallons are purified daily

WATER

The most effective method of destroying organic matter, especially bacteria, is by chemical treatment. Ozone, a gas like oxygen but much more active, is sometimes used (30). In recent years chlorine, or one of its compounds, has come into general use because of its simplicity and certainty. A small amount of liquid chlorine is allowed to flow into the entering supply (Fig. 35). During the war,

Fig. 35. — Plant for the purification of water by chlorine. A chlorinator of the New A'ork City water system by which 400 million gallons are purified daily.

water contaminated through accident or intent was often rendered fit for drinking by means of chlorine compounds.

Water containing bacteria can be purified on a small scale or in an emergency, e.g. for household use during an epidemic, by boiling ten or fifteen minutes, and then putting it in a bottle or jar stoppered with cotton. The bottle or jar for keeping the water should also be boiled in water before use

If the purity of a drinking water is doubtful, a sample should be subjected to a chemical and microscopic examination, supplemented by a rigid sanitary inspection of the surroundings and the source of supply.

PRACTICAL CHEMISTRY

70. Distillation. — Water can be purified on a small or local scale by distillation. This operation is often per- formed in the laboratory in a condenser, which is shown in Fig. 36 arranged for use.

Fig. 36. — Condenser arranged for distillation of water.

The condenser consists of an outer tube AA', provided with an inlet and an outlet for a current of cold water which surrounds the inner tube BB'. The vapor from the w^ater boiling in the flask C condenses in the inner tube, owing to the decrease in temperature, and drops off the lower end of this tube, as the distillate, into the receiver D, while the non-volatile impurities remain behind in the flask. Other forms of condensers are used, especially for continuous work (Fig. 37). Cold water enters the condenser at the lower inlet and is kept level in the chamber by the upper outlet. The chamber is heated, and th-e steam in passing down through the con- denser drops off the lower end as distilled water.

Distilled water is prepared on a large scale

by boiling the water in a metal vessel and

condensing the vapor in a block tin pipe

coiled around the inside of a vessel through which a current of cold

water is flowing (Fig. 38).

Fig. 37. — Apparatus for distilling water

WATER

6,3

Fig. 38. — Coiled pipe condenser

Distilled water has a Hat taste. If it is to be used as a beverage, air is bubbled through the water, which soon acquires the accustomed flavor. i\Iuch distilled water is used in the chemical laboratory to prepare solutions and in experiments requiring water free from organic and mineral matter.

71. Some properties of water. — At ordinary temper- atures pure water is a tasteless and odorless liquid. It is usually colorless, but thick layers are blue. When water is cooled suffi- ciently, the liquid becomes the familiar soKd, ice. The temperature at which water solidifies or freezes is 0° C. (or 32° F. — Fahrenheit). When water freezes, it expands about one tenth of its volume. That is, a given volume of water produces a larger volume of ice. If we weigh the same volume of water and ice, we find the ice weighs less. In other words, the specific gravity of ice is less than i, which is the standard. Since the specific gravity of ice is about 0.92, ice floats in water.

The pressure exerted by water when it freezes is powerful. Vessels 01 pipes completely filled with water often burst when the water freezes. It is a popular idea that " thawing out " a pipe bursts it. This is not true, because ice contracts when it melts. As a matter of fact, pipes crack as soon as the water freezes, and when the ice melts, the water flows out of the crack.

Ice melts at 0° C. (32° F.), which is also the freezing point of water. Ice often crystallizes in forming, but individual crystals are seldom visible except during the first stages of the process. Snow crystals are common (Fig. 30). They are always six-sided or six-pointed, and are formed in the atmosphere by the freezing of water vapor.

PRACTICAL CHEMISTRY

When water is heated sufhciently in an open vessel, the liquid becomes a vapor, and the vapor escapes rapidly

^ until the temperature reaches ioo° C. (or 212° F.). At this point water boils, i.e. it changes rapidly into vapor without rise of

temperature. This Fig. 39. — Snow crystals. (From a photo- . . .

graph by permission of Wilson A. Bentley) vapor IS mviSlble, but

as it leaves the vessel it cools and condenses quickly into a cloud of minute drops of liquid water. This cloud is called steam.

The relation of water vapor to steam can be shown by a simple experiment. Water is put in a flask (Fig. 40) and boiled vigorously. The flask is now full of invisible water vapor and steam is seen es- caping from the bent exit tube. When the steam itself is heated, the cloud disappears from the end of the exit tube, i.e. the vapor no longer condenses into droplets but escapes unchanged as an in- visible gas.

Water exists in three states — liquid, soHd, and gaseous. That is, water, ice, and vapor are three common states of the same compound. And, furthermore, the trans- formation from one state to another is a matter of tem- perature. Hence the tem- peratures at which these ^.^ ^^ _ Experiment to show the transformations take place relation of water vapor to steam

WATER 65

arc important; they should be learned. Change of state does not lead to the formation of a new compound. Such a change is called a physical change.

72. Evaporation. — Licjuid water is always changing into vapor, whatever the temperature. This process is called evaporation. Water vapor is escaping constantly on an enormous scale from the surface of the ocean, lakes, and rivers. This vapor is transported by air currents to great heights and to regions remote from the water surface. When the vapor reaches places where the temperature is low, a part condenses to the liquid or solid state and pro- duces . clouds, rain, snow, or other forms of water. By evaporation heat is absorbed. By condensation heat is lib- erated. These two processes produce a fundamental effect on climate, not only in regulating temperature but in dis- tributing water through the air and upon the land.

On a small scale water vapor is escaping from puddles of water, moist soil, and wet objects. We speak of this process as drying. The rate at which drying takes place depends largely on the amount of water vapor in the air, the rate being in general low on wet days and high on dry days. This subject will be treated in Chapter IX.

73. What is vapor pressure ? — Evaporation produces pressure. This fact is not apparent in the mere escape of water vapor into the air from a surface of water or from a wet object. Nevertheless the pressure is produced and it is called the vapor pressure of water, or vapor pressure.

A simple experiment will illustrate the fact of vapor pressure. The apparatus is shown in Fig. 41. The dry bottle (left) is fitted with an open U-shaped tube partly filled with a colored liquid {e.g. ink) to serve as an indicator of pressure. When the stopper is re- moved, a littlj water poured into the bottle, and the stopper quickly replaced, the water begins to evaporate. As the water evaporates, the colored liquid shows an increase in pressure (right).

PRACTICAL CHEMISTRY

The pressure exerted by water vapor depends solely on the temperature. This is readily seen by comparing the

^£\

' Fig. 41. — Experiment to illustrate vapor pressure

heights of the mercury in the fine-bore tubes shown in Fig. 42. Each tube was first filled with mercury and in- verted in the dish of mercury. In each tube the mercury sank to the same point (760). In tube A there is no water vapor in the space above the mercury, and the height of the mercury column is 760 mm. A small drop of water is forced up into the tubes B and C by means of a medicine dropper. In B the space above the mercury is filled with water vapor at 20° C. ; the vapor exerts a pressure and forces the mercury (^ UK,^ down to nearly 742 mm. That is, water \ / vapor at 20° C. exerts a pressure equal to

18 mm. of mercury. Similarly, in C the

'per^ment to' ^pace is filled with water vapor at 50° C, show the rela- and the mercury is forced down to 668 mm., tion between ^]^g water vapor at this temperature ex- vapor pressure , . r •and tempera- "ting a pressure of 92 mm. ture If the water vapor in a tube were at loo

WATER 67

C, the vapor pressure would be 760 mm. The latter value is instructive, for it means that at the boiling point of water (100° C.) the vapor pressure equals numerically the nor- mal atmospheric pressure. This means in other words that the vapor escaping from water boiUng in an open vessel overcomes the pressure of the atmosphere upon the surface of the water. Since the normal pressure of the atmos- phere is 760 mm., the normal boiling point, so to speak, is 100° C.

The boiling point becomes lower as the pressure is decreased and higher as the pressure is increased. Warm water will boil under the receiver of an air pump and on the top of a high mountain. In a pressure cooker the temperature of the water is above 100° C. In a vacuum vessel, such as is used to evaporate sugar solutions, the boiling point is often as low as 70° C.

The pressure exerted by water vapor depends solely, as shown above, on the temperature of the evaporating water, and has a maximum value for each temperature. These values have been carefully determined by experi- ment, and can be found in the Table of Vapor Pressure given in the Appendix, § 4.

74. A practical application of vapor pressure is made in finding accurately the weight of a Uter of oxygen and in similar experiments where gases are measured over water. The oxygen is collected in a bottle or graduated tube inverted in a vessel of water. The gas should be allowed to stand confined over the water long enough to become saturated with water vapor, i.e. the tube finally contains a mixture of oxygen and the maximum amount of water vapor at the given temperature. In such a mixture, each gas shares the total atmospheric pressure. Hence the actual pressure exerted by the oxygen is found by subtracting the pressure of the water vapor (found in the table) from the total pressure (indicated by the barometer). The corrected pressure is used in the formula for reducing the volume of a gas to its volume at 0° C. and 760 mm. (65).

68 PRACTICAL CHEMISTRY

75. Chemical conduct of water. — We have already seen that water at ordinary temperatures interacts with cer- tain metals, especially calcium, sodium, and potassium, and at higher temperatures with iron (49). Water is de- composed to some extent into its component elements (oxy- gen and hydrogen) by intense heat ; at about 2000° C. the. decomposition is less than 2 per cent. As the tem- perature falls, the elements recombine to form water. In other words, water is a stable substance towards heat.

Water combines directly with many oxides. Thus^ lime, which is calcium oxide, combines directly with water and forms calcium hydroxide ; this chemical change is often attended by considerable heat, as may be seen when mortar is being prepared. Similarly, sulphur dioxide forms the compound sulphurous acid. These chemical changes may be represented thus : —

Calcium Oxide -j- Water = Calcium Hydroxide

(Calcium-Oxygen) (Calcium-Hydrogen-Oxygen)

Sulphur Dioxide + Water = Sulphurous Acid (Sulphur-Oxygen) (Sulphur-Hydrogen-Oxygen)

Such oxides are called anhydrides.

Water combines with certain solids when they separate from a solution by crystallization. Thus, from a solu- tion of copper sulphate blue crystals are obtained, which when heated give off water and crumble to a gray white powder. (See Water of Crystallization, 83.)

76. Solvent power of water. — Water dissolves a great many substances, i.e. they disappear when put into w^ater. The liquid which results from this process of dissolving is called a solution. The dissolved substance is distributed uniformly throughout the whole liquid and will not settle out. The liquid in which the substance dissolves is called

WATER 69

the solvent. The dissolved substance is called the solute. A water solution of a substance is sometimes called an aque- ous solution — the kind we are considering in this chapter. Substances differ widely in their solubility. A solution which contains a small proportion of solute is called a dilute solution ; one containing a large proportion of solute is called a concentrated solution.

77. Solution of gases. — Water dissolves many gases. Some like ammonia, are very soluble, while others, such as oxygen and nitrogen, are only shghtly soluble. Air dis- solves sHghtly in water. Thus, if faucet water, or water that has been shaken in a bottle, is heated, bubbles gather and escape as the heat is increased. As a rule, the solu- bility of a gas decreases with rise of temperature. Pres- sure, too, influences the solubility of gases. Thus, as we have already seen (36), carbon dioxide is forced into cyHnders of water in preparing soda water. W^hen the pres- sure is decreased by opening the valve, the gas escapes rapidly and causes the soda water to froth or foam. Bubbles caused by escaping carbon dioxide may also be seen when the stopper is removed from a bottle containing a charged beverage (Figs. 14, 15).

This rapid escape of a gas is called effervescence. Underground waters often contain considerable carbon dioxide, owing to the great pressure to which subterra- nean gases are subjected. Hence, many natural mineral waters effervesce when they come to the surface.

78. Solutions of liquids. — Some liquids, such as alcohol and glycerin, dissolve in water in all proportions ; others, e.g. gasolene and kerosene, are very slightly soluble, as is shown by the fact that after agitation with water they sep- arate almost entirely as distinct layers of oil and water.

Some Uquids dissolve to a limited extent in water. Ether

yo PRACTICAL CHEMISTRY

is an example. If ether is shaken with water and the liquid is allowed to stand a short time, two layers form. The upper layer consists of ether and a httle water ; the lower layer is the opposite.

On the other hand alcohol and water form no such layers, not simply because each is soluble in the other, but because each is soluble without limit in the other ; that is, alcohol- water is a case of perfect mutual solubility, whereas ether-water is a case of limited solubility.

79. Solutions of solids. — Water dissolves many sohds, and such solutions are very useful. The solubility of solids in water is a matter of the utmost practical im.portance, not only in its far-reaching effect in nature but also in its indispensable use to man.

The solubihty of sohds in water depends on the substance. Some, like sand, dissolve scarcely at all and are often described as insoluble. Others, like salt, are moderately soluble, while still others, like magnesium chloride or sodium hydroxide, are highly soluble.

The degree of solubility depends on temperature. In most cases solubihty increases with rise of temperature ; hence the common practice of heating to hasten solution. A few sohds (e.g. calcium hydroxide) are less soluble in hot water than in cold, and a few others (e.g. sodium chlo- ride) dissolve to about the same degree in hot and cold water.

There is a limit to the solubility of most substances. As a general rule, a given weight of water at a fixed tempera- ture will dissolve only a definite weight of solid ; and this is the case, even though more undissolved sohd is available for solution. A solution conforming to the conditions just stated is said to be saturated. For general purposes, solubility may be expressed by such terms as insoluble,

WATER

Table of the Solubility of Solids ix Water

Solids

Number of Gr.\ms in Solution in 100 Grams of Water

20° C.

100° C.

Calcium chloride Calcium hydroxide Magnesium sulphate Potassium dichromate Potassium, nitrate Sodium chloride

74-5 0.165 36.2 13.0 31.6 36.0

159.0 0.077

73.8 102.0 246.0

39.8

slightly soluble, or very soluble. It is more accurate to represent the amount of solvent by loo gm. ; on this basis the solubihty of a sohd is the number of grams of soHd dis- solved by 100 gm. of water.

80. Solubility tables and curves. — The accompanying table of solubihty shows the solubihty of some solids. But this table is hmited, only two temperatures being given.

A better way of representing the solubihty of a sub- stance is by a solubihty curve. The curves of several substances are shown in Fig. 43. The temperature is read along the vertical lines and the number of grams of solute in 100 gm. of water along the horizontal lines. Many facts can be read from solubihty curves. For example, if we wish to know the temperature at which 40 gm. of potassium chlorate are held in solution by 100 gm. of water, it is only necessary to find where the horizontal hne numbered 40 cuts the potassium chlorate curve, and then follow the vertical line down to the temperature number, where 80° C. is found.

81. Solution and crystallization. — If hot solutions are cooled or if enough of the solvent is removed by evaporation^ the solute separates from the solvent in crystals. The process of separating a dissolved solid from its solution by

PRACTICAL CHEMISTRY

40' 50^ Temperature Fig. 43. — Solubility curves

either of these ways is called crystallization. The shape and color of the crystals are characteristic of the particular substance and serve to identify it. Thus, common salt crystallizes in white cubes. (See also 83.)

82. Supersaturated solutions. — Crystals are not always deposited from a cooled or a concentrated solution, as just stated. Thus, a hot, very concentrated solution of some soHds, such as sodium sulphate or sodium thiosulphate, deposits no crystals when the clear solution cools, although the solution actually contains more solute than the solvent

WATER

could ordinarily dissolve at the lower temperature ! Solu- tions which contain more solute than is needed for normal saturation are called supersaturated. Supersaturation oc- curs only under special conditions. For example, if a super- saturated solution is stirred or vio- lently shaken, crystals begin to form. Moreover, if a fragment of the solid is dropped into the super- saturated solution, crystals very quickly form upon the fragment and soon accumulate in a conspic- uously large quantity (Fig 44).

Fig. 44. — Experiment illus- trating supersaturation

Supersaturation can be shown by a simple experiment. A test tube containing two or three cubic centimeters of water and considerable crystallized sodium thiosulphate is heated until a solution is produced. The clear solution is then poured into a warm test tube, a cork or wad of cotton is inserted, and the solution is allowed to cool. It remains clear. Now if a fragm.ent of a crystal of the solid is dropped in, the crystals soon form and may become a solid mass. In Fig. 44 the successive stages are shown.

83. Solution and water of crystallization. — Crystals of some solids contain water, which is an essential part of the compound. The combined water must not be confused with water which adheres to a crystal or is inclosed in it. Crystals containing combined water are dry even after the crystals are powdered. The combined water can be removed by heat or sometimes merely by exposure to air. Loss of water is usually attended by loss of color and always by loss of crystalHne appearance. Thus, crystallized cop- per sulphate loses its blue color slowly at ordinary tempera- tures and very rapidly when heated, finally becoming a gray powder.

74 PRACTICAL CHEMISTRY

The proportion of combined water in crystals is constant in the same compound, but in different substances the proportion varies between wide hmits. Water chemi- cally combined in a crystal and readily removed in a definite proportion by heating is called water of crystal- lization.

Compounds containing water of crystallization are some- times called hydrates or hydrated compounds. Conversely, compounds which have been deprived of water of crystalli- zation are said to be anhydrous or dehydrated. For ex- ample, blue crystallized copper sulphate is a hydrate of the compound copper sulphate ; but after the blue compound has been heated, it is anhydrous or dehydrated copper sulphate, which is a gray powder. Anhydrous compounds often readily become hydrated again. Thus, when the gray anhydrous copper sulphate is added to water, a blue solution is produced from which blue crystals of hydrated copper sulphate are readily obtained.

84. Efflorescence. — Some crystallized substances lose water of crystallization merely by exposure to air and be- come powders. This property is called efflorescence, and such substances are said to be efflorescent or to effloresce. Crystals of washing soda, alum, and borax effloresce readily.

An explanation of efflorescence is found in the principle of vapor pressure. Solids containing water of crystallization exert a slight vapor pressure. If this vapor pressure is greater than the pressure of the water vapor in the atmosphere, the substance loses water until the vapor pressures are equal or until all the water has escaped from the substance. On the other hand, some substances containing water of crystallization do not effloresce because their tendency to lose water is counteracted by the water vapor in the air. Crystallized barium chloride and gypsum belong to the latter class.

WATER 75

85. Deliquescence. — Many substances when exposed to air become moist, and sometimes even dissolve in the ab- sorbed water. Calcium chloride, potassium carbonate, zinc chloride, sodium hydroxide, and potassium hydroxide belong to this class. This property is called deliquescence and the substances are said to deli- quesce, or to be deliquescent. Deli- quescence is a property of very soluble substances. Common salt often deli- quesces, especially in damp weather, owing to small quantities of magnesium and calcium chlorides which are present

as impurities. The property of deli- / '

quescence is utilized in the laboratory '^^-^:^^:^-^

to dry substances, calcium chloride Fig. 45- — A desiccator

1 . r. 1 J r iU' containing calcium

being often employed for this purpose. ^^,^^.j^ ^.^ ^^^ ^^^_ One form of apparatus used is called torn) a desiccator (Fig. 45) .

Deliquescence is readily explained. Water vapor from the air condenses on the surface of the solid and produces a very concentrated solution, which has a vapor pressure much lower than the average pressure of the water vapor in the air ; the solution, therefore, con- tinues to take up water until its vapor pressure equals the pressure of the water vapor in the air.

86. What is a solution ? — Let us try three experiments : (i) If we shake a httle salt with water, the salt dissolves. We have a solution, i.e. a mixture in which the salt is uni- formly distributed and from which the salt will not settle out. (2) If we shake a Httle powdered starch with cold water, the starch is distributed more or less uniformly throughout the liquid. But on standing, the starch begins to settle out, and finally the mixture separates into water and starch. The starch was not dissolved, but merely

76 PRACTICAL CHEMISTRY

suspended in the water. Such a mixture of a soHd and water is called a suspension. (3) If we shake kerosene oil vigorously with water, the oil breaks up into fine drops which are distributed throughout the liquid. After a time, however, the oil separates from the water. Such a mix- ture of a liquid and water is called an emulsion.

Suspensions and emulsions have a common property, viz. sooner or later the substance separates from the w^ater. On the other hand, the substance does not separate from a solution. This distinctive property of a solution, viz. absence of settling, is doubtless due to the fact that in a solution the dissolved particles are exceedingly minute — too minute to be seen through a microscope or detected by a beam of hght. It w^as found many years ago that if a beam of Hght is passed into a dark room, the path is re- vealed by the dust particles that reflect the light; this effect may be seen when a sunbeam comes through an opening in a blind or a hole in a curtain. But when a strong beam of light is passed through a solution, no bright path is revealed, because the dissolved particles are much too small to reflect light.

Certain substances, however, form clear mixtures with water from which the substance does not settle nor can it be removed by filtering. These liquids look homogeneous, i.e. all parts appear to be just alike. And yet, if we pass a beam of hght through them, the path is bright, thereby proving that these mixtures contain particles in suspension. Such mixtures are not true solutions, but colloidal solutions, i.e. mixtures in which the particles in suspension are very fine — only a little coarser, in fact, than the particles in a true solution. Colloidal solutions should really be called colloidal suspensions, but the term solution is often used. Many substances can be reduced to the colloidal state.

WATER 77

They are called colloids. Typical examples are starch (349), silicic acid (389 1, clay, and metals like gold. (See especially 101.)

In true solutions, then, the particles of the solute are in an exceedingly fine state of division. Whereas in suspensions and emulsions the particles are much larger — large enough to be seen through a microscope and often with the eye. Between these two classes come colloidal solutions ; in them the particles are too fine to settle out, though large enough to reflect light, ranging in size from those in true solutions to those in typical suspensions (101). In passing from true solutions through colloidal solutions to sus- pensions (and emulsions), the change is gradual, not abrupt, since the distinction seems to be based fundamentally on the size of the particles.

87. Composition of water. — Water is a compound of hydrogen and oxygen. That is, its constituents are the elements hydrogen and oxygen, and they are chemically combined in a constant ratio.

We have already learned in several ways that hydrogen and oxygen are constituents of water, (i) Metals, such as calcium, sodium, and iron, liberate hydrogen from water and form simultaneously com- pounds containing oxygen (49). (2) If an electric current is passed through an acid (or alkaline) solution of water, hydrogen and oxygen are liberated (16). (3) When hydrogen is burned iA air or in oxygen, water is produced (53).

We can show by a simple experiment that oxygen is a chemical constituent of water. A tube about a meter long and closed at one end is completely filled with chlorine water (prepared by saturating water with chlorine — an element to be studied in Chapter XI). The open end is immersed in a vessel containing some of the same solution, and the whole apparatus is placed in the direct sunlight.

PRACTICAL CHEMISTRY

The chlorine and water interact, forming hydrochloric acid and oxygen. Bubbles of gas soon appear in the liquid, and after a few hours a small volume of gas collects at the top of the tube (Fig. 46). The gas can be shown to be oxygen by the usual test, viz. relighting a glowing joss stick or splint of wood.

88. Electrolysis of water. — The decom- position of water by electricity, or, as it is called traditionally, the electrolysis of water, shows by a single experiment that water con- sists of the elements hydrogen and oxygen. It is done in the laboratory in a special form of apparatus (Fig. 47).

not conduct electricity, so a 0 vols.) and

Fig. 46. — Pre- paring oxygen from water by chlorine

Pure water does Y~^"A^ J~\ mixture of water (

concentrated sulphuric acid (i vol.) is poured into the apparatus until the reservoir is half fidl (after the stopcocks have been closed) . As soon as an electric battery of three or more cells (or a reduced street current) is connected with the piece of platinum near the bottom of each tube, bubbles of gas appear on the platinum, rise, collect in the upper part of the tubes, and slowly force liquid from each tube into the reservoir.

The volume of gas is greater in one tube than in the other, when the electrolysis is stopped. As- suming that the tubes have the same diameter, the gas volumes are in the same ratio as their heights, which will be found by measurement to be ap- proximately two to one. Tests applied to each gas (by letting a little out through the stopcock) show that the gas having the larger volume is hydrogen pig. 47.

and that the other gas is oxygen.

Elec- trolysis of water

WATER

89. How the exact composition of water is found. — The experiments just cited and described show the qualitative composition of water. That is, they show that water is a compound of the two elements hydrogen and oxygen. But they give us no information about the propor- tion of the elements in the compound. To find the quan- titative composition of water, we must study the results of experiments per- formed for the purpose of determining the exact pro- portions — " the quan- tity " — in which the two elements combine to form the compound.

Since the constituents of water are gases, we can find the composition by volume, i.e. volume tri- cally, as well as by weight, i.e. gravimetrically. More- over, we can determine the quantitative composition by analysis, i.e. taking the compound apart chemically, or by synthesis, i.e. putting its parts together chemically.

90. Morley's determination of the gravimetric compo- sition of water by synthesis. — The most accurate deter- mination of the gravimetric composition of water ever made was completed by the American chemist Morley in 1895 (Fig. 48). In his experiments he not only weighed the hydrogen and oxygen that combined but also the water formed by their synthesis.

The apparatus (Fig. 49) was weighed vacuous {i.e. free

Fig. 48. — Morley (1838 )

PRACTICAL CHEMISTRY

from air or other gases). The tubes aa were connected with the weighed reservoirs of pure oxygen and hydrogen, and the oxygen was introduced. Sparks were next passed between the platinum wires cc, and the heat ignited the hydrogen, which was slowly ad- mitted, the combination of gases taking place at bb. The water vapor condensed in the tube dd, the lower portion of which was immersed in cold water. The combustion of the hydro- gen was continued until a suitable weight of water was formed. The introduction of the gases was then stopped. The water and its vapor were then converted into ice by putting the apparatus into a freezing mixture. The gases left over when the combustion was stopped were drawn off, passing in their exit through tubes of phosphorus pentoxide in ee which re- tained all traces of water. This mixture was M or ley's analyzed and allowance made for the hydrogen apparatus ^^ oxyffcn in it. The whole apparatus, freed

for deter- , , ^, , r n • -u j

mining the ^ ^om hydrogen and oxygen, was finally weighed ; gravimet- the increase was the weight of water formed by ric com- ^Yie combination of known weights of hydrogen water^" ° and oxygen. As the result of exceptionally ac- curate experiments Morley found that i part by weight of hydrogen combines with 7.9395 parts by weight of oxygen.

91. Dumas' determination. — Another method was used by the famous French chemist Dumas in 1843. It consisted in passing dry hydrogen over heated copper oxide. We have seen (55) that in this chemical change hydrogen reduces copper oxide and thereby forms copper and water. If the copper oxide and copper are weighed, the loss is the weight of the oxygen used. If the water is collected and weighed, the difference between the weights of the water formed and

Fig. 49.

WATER

the oxygen used is the weight of the hydrogen. The result obtained by Dumas has long ceased to be regarded as accurate, but the de- termination has historical interest.

92. The volumetric composition of water. — The electrol- ysis of water is good evidence that the volumetric compo- sition of water is 2 of hydrogen to i of oxy- gen. An accurate determination is made by synthesis, i.e. by measuring accurately the volumes of hydrogen and oxygen that com- bine to form water. A sketch of one form of apparatus is shown in Fig. 50. The essential part is the eudiometer A. It is a graduated glass tube closed at the upper end. Near this end two platinum wires are fused into the glass ; the outer ends are looped and the inner ends are near together so that an electric spark will leap across the gap and produce enough heat to cause the two gases to unite.

Fig. 50. — Ap- paratus for determining the volumet- ric compo- sition of water

The experiment is readily performed. The eudiom- eter is filled with mercury and inverted in the jar of the same liquid. Hydrogen is introduced until the eudiometer is about one fourth full. The mer- cury levels inside and outside are made the same, and the volume of hydrogen is read accurately. The temperature and pressure are also read. Approximately an equal volume of oxygen is introduced, the levels are adjusted, and the total volume read accu- rately. Each volume is corrected for temperature and pressure. The difference between the two corrected volumes is the volume of oxygen. An excess of oxygen is needed to lessen the violence of the explosion ; this excess takes no part in the chemical change.

The combination of the two gases is caused by connecting the looped ends of the platinum wires with an induction coil and battery, and .passing a spark across the gap. A slight explosion indicates combination. The mercury, after the shock from the explosion, rises and nearly fills the eudiometer. The volume of water is too

PRACTICAL CHEMISTRY

(>=c^

minute to measure in this apparatus. After the mercury and re- sidual gas (oxygen) are cool, the levels are adjusted, and the volume of gas is read (as well as the temperature and pressure). The corrected volume of this gas is subtracted from the volume of oxygen originally introduced, thus giving the actual volume of oxygen that combined with all the hydrogen to form water.

An example will make this experi- ment clear. Suppose the corrected volumes were : —

Volume of hydrogen added . . 22.3 cc.

Volume of hydrogen and oxygen

added 41.5 cc.

Volume of oxygen added . , 19.2 cc.

Volume of oxygen left ... 8.0 cc.

This means that 19.2 — 8, i.e. 11. 2 cc. of oxygen were actually used. In other words the two gases com- bined in the ratio of 22.3 to 11. 2, or very nearly 2 volumes of hydrogen to I volume of oxygen.

Very exact experiments show that the volumetric ratio of hydrogen to oxygen in water is 2.00268 to i.

93. Gay-Lussac's law of gas vol- umes. — A modification of the ex- periment described in 92 shows that 2 volumes of water vapor are formed when 2 volumes of hydrogen and I volume of ox}^gen unite. In the apparatus (Fig. 51) the eudiometer is surrounded by a large tube through which steam is passed, thereby preventing the condensation of the water vapor. If 2 volumes of hy-

Fig. 51. — Apparatus for showing that 2 vol- umes of water vapor are formed by explod- ing 2 volumes of hy- drogen and I volume of oxygen

WATER

drogen and i volume of oxygen are exploded, 2 volumes of water vapor are formed — provided all the gases are meas- ured at the same pressure and the same temperature (about 100° C). This result means that the volumes of hydrogen, oxygen, and steam are expressed by 2, i, and 2, i.e. by small whole numbers.

When other chemical changes involving gases are studied, similar simple relations are found. This general fact, which was discovered by the French chemist Gay- Lussac (Fig. 52), in 1808, may be stated as a law, thus : —

In a chemical change the volumes of the gases can he expressed by s?nall whole numbers.

This law will be used in a later chapter (XV ; see also 157, 176).

94. Summary of the composition of water. — Experiments show that water consists of the two elements hydrogen and oxygen combined in a fixed ratio by weight, viz. i to 7.9395 ; they are also combined in the ratio of 2.00268 to i by volume. Usually these ratios are stated approximately as 2 to 16 by weight and 2 to i by volume. Often the gravi- metric composition of water is stated inper^cent, viz. 11. 18 per cent of hydrogen and 88.82 per cent of oxygen. Some- times, we say, briefly, water is ^ hydrogen and f oxygen.

95. Hydrogen peroxide. — Hydrogen peroxide is a com-

Gay-Lussac (i 778-1850)

84 PRACTICAL CHEMISTRY

pound of hydrogen and oxygen. But the proportion of the constituents is not the same as in water. Water consists of 1 1. 1 8 per cent of hydrogen and 88.82 per cent of oxygen, whereas hydrogen peroxide consists of 5.88 per cent of hydrogen and 94.11 per cent of oxygen. Stated diflerentiy, water consists of 2 parts of hydrogen and 16 parts of oxygen, and hydrogen peroxide of 2 parts of hydrogen and 32 parts of oxygen — all parts by weight. If we compare the comx- position of water and hydrogen peroxide, we see that in hydrogen peroxide there is twice as much oxygen as in water. This fact is brought out in the old name hydrogen dioxide, which is sometimes used now.

The two compounds, water and hydrogen peroxide, are a good example of the law of multiple proportions (47).

Hydrogen peroxide is a colorless Hquid which dissolves readily in water. It is sold usually as a 3 per cent solution, which has a rather sharp odor and a metallic taste. The hydrogen peroxide decomposes slowly into oxygen and water. Hence it is an oxidizing agent. It is used to bleach {i.e. whiten) hair, feathers, fur, silk, ivory, and bone. It destroys bacteria and is used as an antiseptic, e.g. in cleans- ing wounds, though its effectiveness is often overestimated.

EXERCISES

1. Prepare a summary of (a) the properties of water and (b) the composition of water.

2. Topics for home study, (a) Purification of drinking water. (b) Distillation, (c) Water as an erosive agent, (d) Crystals, (e) Vapor pressure. (/) Chemical conduct of water.

3. Define and illustrate (a) water of crystallization, (b) efflores- cence, (c) deliquescence, (d) anhydrous, (e) dehydrated.

4. Define and illustrate (a) solution, (b) solvent, (c) solute, (d) di- lute, (e) concentrated, (/) unsaturated solution, (g) saturated solution, (h) supersaturated solution, (i) solubility, (j) solubility curve.

WATER 85

6. Practical topics, (i) How would you prove that a given liquid is pure water? (2) What conditions are favorable for (a) evaporation, (b) efflorescence, (c) deliquescence? (3) Suggest experiments (a) to find the solubility of a solid in water at 40° C, {b) to show that water from a crystal is not water of crystallization, (c) to find the per cent of water in a potato. (4) Why is sea water salt? How might sea water be rendered suitable for drinking?

6. State {a) the volumetric and (b) the gravimetric composition of water. How is each found?

7. What do these show about the composition of water: (a) La- voisier's experiment (49)? (b) Burning of hydrogen? (c) Elec- trolysis of water?

8. Describe ]\Iorley's experiment on the composition of water,

9. State and illustrate Gay-Lussac's law.

PROBLEMS

1. Find the volume of the dry gas at 0° C. and 760 mm. in : (a) 80 cc. at 760 mm. and 17° C. ; (b) 80 cc. at 745 mm. and 19° C. ; (c) 100 cc. at 765 mm. and 17.5° C. ; (d) 97 cc. at 757 mm. and 20.5° C.

2. Plot the following data on cross section paper and draw the solu- bility curve of the substance : Temperature — o, 10, 20, 30, 40, 50, 55 ; corresponding solubility (i.e. number of gm. soluble in 100 gm. of water) — 13, 21, 31, 45, 64, 86, 100.

3. If the density of ice is 0.92, what volume will a liter of water at 4° C. occupy w^hen frozen? Ans. 1.087 1-

4. By the use of the solubility curves in Fig. 43 answer the follow- ing : (a) How many gm. of sodium chloride are in solution at 20°, 30°, 55°, 65°, 0°, 100° ? {b) At what temperatures are 60 gm. and 95 gm. of potassium bromide in solution? (r) Compare the solu- bility of sodium nitrate and sodium chloride. How much of each is in solution at 20'', 25°, 30°?

5. Calculate the per cent of water of crystallization in each crystal- lized substance from the following : (a) 5 gm. of aluminium sulphate lose 2.43 gm. on heating ; (b) 7 gm. of calcium sulphate lose 1.464 gm. ; (c) 3 gm. of cadmium nitrate lose 0.7 gm. ; (d) 3 gm. of cobalt nitrate lose 1. 113 gm.

6. Suppose 15 gm. of water are decomposed. What weight of (a) oxygen and (b) hydrogen is produced? What volume (at 0° C. and 760 mm.) of (c) oxygen and (d) hydrogen?

7. What volume of oxygen (at 0° C. and 760 mm.) must be used to unite with 175 gm. of hydrogen to form water?

86 PRACTICAL CHEMISTRY

8. What volume of hydrogen (at o° C. and 760 mm.) must be used to convert 175 gm, of oxygen into water?

9. A mixture of 500 cc. of oxygen and 1250 cc. of hydrogen (both at 0° C. and 760 mm.) is exploded. What weight of water is formed? Ans. 0.805 S'^-

10. 50 cc. of oxygen are mixed with 500 cc. of hydrogen, both measured at the normal temperature and pressure. x\n electric spark is passed through the mixture. What volume, if any, of gas will re- main, and how would you ascertain whether it is hydrogen or oxygen?

CHAPTER VII

LAW AND THEORY — LAWS OF DEFINITE AND MULTIPLE PROPORTIONS — ATOMIC THEORY — ATOMS AND MOLE- CULES-ATOMIC WEIGHTS

96. Facts, laws, and theories. — The facts we study in chemistry were discovered by observation and experiment. Facts which always occur under the same conditions soon become well estabhshed. Many related facts are often summarized in a brief general statement called a law. The explanation we give of facts, especially groups of related facts, is called a theory.

Laws and theories are of great service in chemistry, since they help us gather into inteUigible statements our knowl- edge of a vast number of related facts. They also help us discover new facts and interpret phenomena.

In this chapter wq shall review^ three laws already studied, and interpret them by a theory called the atomic theory.

97. Law of the conservation of matter. — We have seen (5) that in a chemical change the total weight of the mat- ter involved is not altered. That is, the sum of the weights of the substances entering into a chemical change always equals the sum of the weights of the substances resulting from the chemical change. This general fact about chemi- cal change is summed up by the law of the conservation of matter, thus : —

No iveight is lost or gained in a chemical change. Sometimes this law is stated in another form, viz. mat- ter is indestructible.

88 PRACTICAL CHEMISTRY

98. Law of constant composition. — We learned in 9 that constant composition is an essential characteristic of a compound, e.g. water contains 88.82 per cent oxygen and 1 1. 18 per cent hydrogen. Experiments show that in all chemical compounds the different constituents are present in a definite and constant proportion by weight. This general fact, stated in the form of a law, becomes the law of constant composition or the law of definite proportions, thus : —

A given chemical compound always contains the same ele- ments in the same proportion by weight.

99. Law of multiple proportions. — We saw in 47 that if the composition of different compounds of the same ele- ments is expressed in a special way, multiple relations exist between the weights of one constituent. Composition is usually expressed in per cent. But if a fixed weight of one constituent is adopted as a basis, and the composition of the compounds is expressed in terms of this weight, then the simple multiple relation which exists between the weights of the other constituent (or constituents) may be clearly seen.

Consider this case. No multiple relation appears in the statement that the two compounds of carbon and oxygen consist respectively of 27.27 and 42.85 per cent of carbon and 72.72 and 57.14 per cent of oxygen. However, if we state these proportions in a different way, multiple re- lations will be clearly seen. We first reduce each propor- tion of carbon to i thus: 27.27 -^ 27.27 = i, and 42.85 -^42.85 = I. But in order to maintain the correct relation of carbon to oxygen (found by analysis and expressed above as per cent) we must divide each proportion of oxygen by the corresponding number, i.e. 72.72 -^ 27.27 = 2.6, and 57.14 -r- 42.85 = 1.3. These two operations give us: —

LAW AND THEORY

Ratio of carbon i to i Ratio of oxygen 2.6 to 1.3

Now 2.6 and 1.3 are in the ratio of 2 to i. Obviously the weights of oxygen in the carbon oxides are in a simple multiple relation.

The general fact of multiple proportions is expressed as the law of multiple proportions, thus : —

When elements unite to form a series of compounds, a fixed weight of one element always combines with such weights of the other element {or elements) that the ratio between these diferent weights can be expressed by sm.all whole numbers.

100. The atomic the- ory. — The theory that explains the facts sum- marized in the three fun- damental laws just re- viewed is called the atomic theory. It was proposed about 1805 by the English chemist Dal- ton (Fig. 53). According to this theory: (i) an element is made up of a vast number of very small particles called atoms; (2) atoms of the same element have the same weight ; (3) atoms of different elements dij6fer from one another in weight; (4) chemical change is the union, separation, or exchange of whole atoms. The atomic the- ory means in a few words that matter is composed of atoms, which remain undivided in chemical changes.

Fig. 53. — Dalton (i 766-1844)

90 PRACTICAL CHEMISTRY

With the help of the atomic theory, many facts about elements, compounds, and chemical change can be made very much clearer.

101. Atoms and molecules. — An atom is the smallest, indivisible particle of an element. In a chemical change the atom is the vehicle. It is the smallest part of an ele- ment which participates in a chemical change. If two or more atoms are held together by chemical attraction, the product of this combination of atoms is called a mole- cule. If the atoms are alike, the molecule is a molecule of an element, e.g. a molecule of oxygen consists of two atoms of the element oxygen (O2). If the atoms are not alike, the molecule is a molecule of a compound, e.g. a molecule of water consists of two atoms of hydrogen and one of oxy- gen (H2O). A molecule of a compound always contains one or more atoms of different elements ; hence, we may think of such molecules as the smallest particles of the compound. Thus, a molecule of water is the smallest particle of the compound water because if we try to pro- duce a smaller particle, we obtain atoms of the elements hydrogen and oxygen.

Although atoms and molecules are exceedingly small, too small to be seen by direct observation, nevertheless certain experiments give conclusive evidence of their exist- ence. Atoms will be considered later ; we can consider molecules now.

We found in 66 that certain properties of gases can be readily and adequately explained by the kinetic-molecular theory, i.e. the supposition that gases consist of exceedingly minute, rapidly moving particles called molecules.

Are there molecules in a liquid? When a colloidal solu- tion (86) is examined with an ordinary microscope under favorable conditions, tiny moving points of light are seen.

LAW AND THEORY

If an ultra-microscope is used, the result is startling. The tiny points of light seen in the liquid dart about with a zig- zag motion. An ultra-microscope is one arranged so we can look down against a dark background into the liquid through which a beam of strong light is sent horizontally. These moving points are colloidal particles, which, although invisible, reveal their presence by re- flected light.

The path traversed by one of these points seen in a colloidal solution is com- pHcated (Fig. 54). This motion, which is called the Brownian movement, is seen in all colloidal solutions regardless of the nature of the particles (though seen best with finest particles) or the age of the solution. The erratic movement of the colloidal particles is due to the ceaseless, irregular bombardment of the particles by the molecules of water. The movements and the size of the colloidal particles have been carefully investi- gated, and the results furnish conclusive evidence of the existence of molecules in solutions.

102. Interpretation of chemical change by the atomic theory. — Let us picture a chemical change in terms of the atomic theory, e.g. the combination of copper and oxygen. A piece of the element copper consists of many miUions of atoms of copper ; a mass of oxygen likewise consists of a very large number of atoms of oxygen. When the chem- ical change occurs between copper and oxygen, atoms of copper combine with atoms of oxygen and form molecules of the compound copper oxide. And this combining of atoms into molecules continues until all the atoms of cop-

Fig. 54. — Path of a colloidal particle

92 PRACTICAL CHEMISTRY

per or all the atoms of oxygen (or under certain conditions all the atoms of both substances) have been used up. Sim- ilarly, the decomposition of mercury oxide is the sepa- ration of numberless molecules of the compound mercury oxide into atoms of the elements mercury and oxygen. So also, the liberation of hydrogen from sulphuric acid is the exchange of atoms of zinc for the atom.s of hydrogen in the compound sulphuric acid.

103. How the atomic theory assists us in explaining the three fundamental laws of chemical change. — These laws, we have already found, are the law of the conser- vation of matter, the law of constant composition, and the law of multiple proportions.

Let us take first the law of the conservation of matter, viz. unvarying weight in a chemical change. According to the atomic theory the weight of an atom is never changed. In the case of copper and oxygen and all other chemical changes, no atoms are created or destroyed ; they are merely redistributed on another plan. Nor are atoms changed in weight; the weight of the copper oxide formed equals the sum of the weights of the copper and oxygen used up. Inasmuch as all chemical changes have this character- istic, viz. unvarying total weight, it is obvious that the atomic theory, which assumes unchanging weights of atoms, is in accord with the law of the conservation of matter.

Second, we consider the law of constant composition, viz. a given compound always consists of the same elements in a fixed proportion. According to the atomic theory, molecules are formed by the union of some whole number of atoms of an element with some whole number of atoms of another element or elements. Each molecule of the com- pound copper oxide, for example, would therefore consist of one or more atoms of copper united with one or more

LAW AND THEORY 93

atoms of ox>^gen, and the composition of each molecule of copper oxide would be constant, i.e. each molecule would consist of the same elements united in a constant ratio by weight. This means that the composition of copper oxide would always be a certain per cent of copper and a certain per cent of oxygen. Since all other chemical compounds have been found to have a constant composition, the atomic theory is in harmony with the law of constant composition.

Finally, we interpret the law of multiple proportions, viz. simple multiple relations exist between the weights of one of the constituents of a series of compounds. Ac- cording to the atomic theory atoms are transferred as wholes ; this means there are no fractions of atoms. Con- sider the two oxides of carbon. Each contains carbon and oxygen in a definite ratio, but the ratios are different. In one the ratio of carbon to oxygen is i to 1.33, and in the other I to 2.66 (99). That is, there is twice as much oxy- gen combined with carbon in one case as in the other. In other words, if a molecule of one compound consists of one atom of carbon and one of ox}^gen, a molecule of the other consists of one of carbon and two of oxygen. Since other series of compounds exhibit a simple multiple rela- tion, the atomic theory agrees with the law of multiple proportions.

104. Atomic v/eights. — According to the atomic theory atoms of the same element always have the same weight but atoms of different elements have different weights. This means (i) that an atom of oxygen, for example, throughout all its varied changes retains its weight, and (2) that this weight differs from the weight of other kinds of atoms. The weights of different kinds of atoms are called the atomic weights of the elements or briefly atomic weights. These weights have been determined by very accurate

94 PRACTICAL CHEMISTRY

experiment. A table giving the exact, as well as the ap- proximate, values can be found on the inside of the back cover of this book.

Atoms are exceedingly small, so small, in fact, that no attempt is made to weigh a single atom. If, however, the proportions in which elements combine are reduced to a special standard, we obtain the relative weights of atoms. The standard atomic weight is oxygen = i6. The atomic weight of carbon is 12, of hydrogen i, of copper 63.57. Other weights can be found in the table.

The approximate atomic weights are accurate enough for general use, and the common ones should be learned.

It must be remembered that the atomic weights are rela- tive weights. That is, the atomic weight of copper is 63.57, not 63.57 gm. or any other actual weight, but 63.57 ^s long as 16 is accepted as the standard atomic weight of oxygen.

The exact determination of atomic weights is a difficult task. Several principles must be considered in making the final selection. Until this subject is discussed (see Chapter XVI), it will be well enough to regard atomic weights as the numerical values of the ele- ments in chemical changes and to select the approximate weights from the table as needed.

EXERCISES

1. Define law and theory as used in science.

2. State the law of conservation of matter. Illustrate it by the de- composition of mercury oxide.

3. State the law of constant composition. Illustrate it by the gravimetric composition of water.

4. State the law of multiple proportions. Illustrate it by the two carbon oxides.

5. State the four points in the atomic theory.

6. (a) What is an atom? A molecule? (b) Discuss the relation of atoms to molecules.

7. Describe a chemical change in terms of the atomic theory.

LAW AND THEORY 95

8. Interpret by the atomic theory the three laws: (a) conserva- tion of matter, (b) constant composition, (c) multiple proportions.

9. What are atomic weights?

10. Learn the approximate atomic weight of these elements : Hydro- gen, oxygen, nitrogen, carbon, copper, iron, sulphur, chlorine.

11. As in E.xercise 10: Potassium, sodium, calcium, lead, mag- nesium, mercury, silver, zinc.

CHAPTER VIII

SYMBOLS AND FORMULAS

105. What do symbols mean? — We learned in 12 that each element is designated by a symbol. Thus, O is the symbol of oxygen, C of carbon, H of hydrogen, Cu of copper.

The symbols of the ele- ments are given in the table on the back inside cover. These symbols were introduced and first used extensively by the Swedish chemist Berzelius (Fig. 55). He analyzed many sub- stances and used the symbols to express the explicit relations of ele- ments and compounds in chemical changes.

Symbols are not merely abbreviations of the names of the ele- ments. Each symbol stands for one atom of an ele- ment. Thus, H represents one atom of hydrogen. If more than one free atom is to be designated, the proper numeral is placed before the symbol. Thus, 2H means 2 free atoms of hydrogen. If we wish to represent atoms in chemical combination, either with themselves or with

Fig. 5

lius (i 779-1848)

SYMBOLS AND FORMULAS 97

other atoms, a subscript is used instead of a coefficient. Thus, H2 means 2 combined atoms of hydrogen, as in H2

or H2O.

Symbols not only represent atoms, but they also express atomic weights. Thus O represents an atom of oxygen which has the atomic weight 16.

To sum up, a symbol has these meanings: (i) an ele- ment, (2) one atom of the element, (3) the atomic weight of the element.

106. What are chemical formulas? We learned in 13 that a formula is a group of symbols which expresses the composition of a compound. In writing a formula, the symbols of the atoms making up a molecule of the com- pound are placed side by side. Thus, CO is the formula of carbon monoxide, because one molecule of this compound consists of one atom each of carbon and oxygen. Whereas CO2 is the formula of carbon dioxide, because one mole- cule is composed of i atom of carbon and 2 atoms of oxygen. The symbols making up a formula might be written in different orders, but usage has determined the order in most cases.

A formula represents one molecule. If we wish to des- ignate several molecules, we place the proper numeral before the formula. Thus, KCIO3 means i molecule and 2KCIO3 means 2 molecules of potassium chlorate. Sim- ilarly, 2O2 means 2 molecules of oxygen each contain- ing 2 atoms.

In certain compounds some of the atoms in a molecule act chemically like a single atom. This fact is ex-pressed by inclosing the group in a parenthesis, e.g. calcium hy- droxide has the formula Ca(0H)2 because the group OH often acts like an atom. Sometimes the parenthesis is replaced by a period, e.g. C2H5.OH (ethyl hydroxide) and

98 PRACTICAL CHEMISTRY

CUSO4.5H2O (copper sulphate pentahydrate (83)). The period and parenthesis are omitted if the composition of the compound is well understood, e.g. ammonium hy- droxide, NH4OH.

If a group of atoms is to be multiplied, it is placed within a parenthesis. Thus, the formula of lead nitrate is Pb(N03)2. This means that the group NO3 is to be multi- plied by 2. The expression 2Pb(N03)2 means that the whole formula must be multiplied by 2. That is, in two molecules of lead nitrate there are 2 atoms of lead, 4 of nitrogen, and 12 of oxygen.

Formulas are the outcome of experiments. We shall learn more about formulas.

107. What do we mean by molecular weights? — Since a symbol stands for the atomic weight of an element, a formula stands for the sum of the atomic weights repre- sented by the group of symbols. This sum is called a molec- ular weight. In a few words, a symbol stands for an atomic weight and a formula stands for a molecular weight. Thus, the symbols H and CI stand for the atomic weights

1 and 35.5 respectively, and the formula HCl stands for their sum i -f 35.5, or 36.5. If we know the formula of a compound, a simple way of finding the molecular weights is to add the atomic weights corresponding to the atoms in the formula. Using approximate values, the molecular weight of water (HoO) is 2 + 16 = 18. Similarly, the molecular weight of lead nitrate (Pb(N03)2) is 207 +

2 (14 + 48) = 33^'

108. How a formula expresses composition. — We have seen by several examples that the composition of a com- pound can be expressed in per cent and also by a formula. Thus, we can express the composition of w^ater by the formula H2O, and also by hydrogen = 11. 18 per cent and

SYMBOLS AND FORMULAS 99

oxygen = 88.82 per cent. How are these chemical and mathematical expressions related? The answer is simple. The symbols stand for numbers, and the mathematical ex- pression is the equivalent of the chemical formula.

Suppose we have the formula of potassium chlorate, KCIO3, and we wish to express the composition in per cent. The process consists in transposing the chemical formula KCIO3 into the equivalent mathematical expression. This is done by : (i) dividing the weight of each element by the molecular weight and (2) multiplying the quotient by 100. Let us take an example. The formula KCIO3 stands for K = 39, CI = 35.5, and 3O = 48 {i.e. 3 X 16), or the molecular weight is 122.5. Then: —

0.3184, or 31.84 per cent of potassium

39 122.5

'^^'^ = 0.2898, or 28.98 per cent of chlorine 122.5

122.5

= 0.3918, or 39.18 per cent of oxygen

Total 100.00

The process of finding the composition in per cent from the formula is often called calculating the percentage com- position of a compound.

109. The simplest formula of a compound can be calcu- lated from its percentage composition. — The calculation of a formula from the percentage composition is simply the process of finding the small whole numbers by which each atomic weight, as represented by its symbol, must be multi- plied in order to express the composition. The process is as follows: (i) divide each per cent by the corresponding atomic weight ; (2) reduce the quotients to whole numbers ; (3) write the formula. Let us take an example. The com- position of sulphuric acid is hydrogen = 2.04 per cent,

lOO PRACTICAL CHEMISTRY

sulphur = 32.65 per cent, oxygen = 65.31 per cent. If the percentage of each element is divided by the corre- sponding atomic weight, the quotients are 2.04, 1.02, and 4.08. Reducing these quotients to whole numbers (by dividing by 1.02 in this case), the final quotients are 2, i, 4. These quotients represent the ratio of the atomic weights in a molecule of this compound. And since atoms are rep- resented by symbols, a molecule of sulphuric acid contains 2 atoms of hydrogen, i of sulphur, and 4 of oxygen. This means that the formula of sulphuric acid must be H2SO4. Formulas calculated by this method are called simplest formulas. (See Determination of Molecular Formulas of Compounds, Chapter XV.)

EXERCISES

1. Prepare a summary of this chapter (a) in the form of short topics and (b) as a connected narrative.

2. What three meanings do symbols have?

3. Learn the symbols of these elements: (a) Aluminium, argon, barium, bromine, calcium, carbon, chlorine, copper, fluorine, gold. (b) Helium, hydrogen, iodine, iron, lead, magnesium, manganese, mercury, nitrogen, oxygen, (c) Phosphorus, potassium, silicon, silver, sodium, sulphur, tin, zinc.

4. What elements correspond to C, CI, Ca, Cu, S, Si, Sn, ]SIg, Mn, Hg, H, He, A, Al?

5. What do these mean? (a) H, Ho, 2H, 2O, O2, 2O2, CI, CI2, 3CI, 3CI2; (b) N2, K, 2Ca, 3Fe, So, Cu, Alo.

6. What is a formula? Illustrate. What does a single formula represent?

7. State all that these mean. (a) H2O, 2H0O, KXO3. 4H0SO4, NaOH, NH4OH, CH3.OH, 3Ca(OH)2, Al2(S04)3, HNO3; (b) BaCl2.2HoO, 2FeS, 3CaCl2, C12H22O11, 2ZnCl2, 3A1(0H)3, 4Ba(N03)2.

8. Give from memory the formulas of the following compounds : water, potassium chlorate, sulphuric acid, magnesium oxide, copper oxide, sodium hydroxide.

9. How many atoms of the different elements are in each formula in Exercise 7?

SYMBOLS AND FORMULAS loi

10. Define and illustrate molecular weight. How is it related to (a) a formula and (b) symbols?

11. What is the molecular weight of each substance in Exercise 7?

12. How does a formula express composition ? Illustrate by KNO3.

13. State the steps in calculating the simplest formula of a com- pound.

14. As in Exercise 13, in calculating percentage composition.

PROBLEMS

1. Calculate the molecular weight (or multimolecular weight) of the following compounds by finding the sum of the atomic weights :

(a) magnesium oxide (]\IgO), (b) hydrogen peroxide, (c) zinc chlo- ride (ZnCl-i), (d) 2Cu(N03)2, (e) 3.^12(804)3, (/) potassium ferrocya- nide (K4Fe(CN)6), (g) 2Na2B407, {h) crystallized ferrous sulphate (FeS04.7H20).

2. Calculate the simplest formula of the compounds which have the following percentage composition: (a) CI = 60.68, Xa = 39.31;

(b) S = 23.52, Ca = 29.41, O = 47-05; (c) C = 40, H = 6.67, O =

53-33-

3. As in Problem 2: (a) N = 26.17, H = 7.48, Ci = 66.35; (b) As = 75.8, O = 24.2; (c) N = 82.35, H = 17.65.

4. As in Problem 2 : (a) Si = 19.5, C = 66.62, H = 13.88 ; (b) Ca = 38.71, P = 20, O = 41.29; (;:) H = I, K = 39.06, C = 11.99, O =

47-95-

5. Calculate the formula of a compound 18 gm. of which contain 8.4 gm. of iron and 9.6 gm. of sulphur.

6. As in Problem 5: 0.84 gm. contain 0.587 gm. of iron and 0.253 gm. of oxygen.

7. Calculate the percentage composition of (a) hydrochloric acid, (b) hydrogen sulphide (H2S), (c) ammonia (XH3), (</) hydrogen peroxide.

8. As in Problem 7 : (a) calcium oxide (CaO), (b) calcium carbonate (CaCOs), U) calcium sulphate (CaS04), (d) calcium fluoride (CaF2).

9. As in Problem 7 • (a) cane sugar (C12H22O11) and (b) grape sugar (CeHriOe).

10. As in Problem 7: (a) sodium phosphate (Xa3P04), (b) diso- dium phosphate (HXa2P04), (c) monosodium phosphate (H2XaP04), (d) phosphoric acid (H3PO4).

11. Calculate the per cent of (a) copper and (b) water in crystallized copper sulphate (CUSO4.5H2O).

12. Calculate the per cent of (a) F in SiF4, (6) Al in AIPO4, {c) O in Mn02, id) Pb in PbCOa-

CHAPTER IX

NITROGEN — THE AIR — ARGON AND HELIUM — LIQUID AIR

110. Introduction. — Reference has already been made to the element nitrogen and its presence in the air. We have also learned a httle about the atmosphere — the great mass of gas that envelops the earth and extends several miles into space. The terms atmosphere, the air, and air are often used interchangeably ; though by the air or air we usually mean a Umited portion of the atmosphere, e.g. the air over a city, the air of a room, or a bottle of air.

In Chapter II we studied oxygen, and learned among other things that this very active gas makes up about one fifth of the air. Since most of the remainder is nitro- gen, we shall study this element before considering the air.

111. Occurrence of nitrogen. — The elementary gas nitrogen constitutes about four fifths (or exactly 78.122 per cent) of the air. Nitrogen is a constituent of nitric acid (HNO3) and ammonia (NH3), and of important com- pounds related to them. Nitrogen is also found in many animal and vegetable substances essential to life, e.g. the compounds called proteins, which are indispensable in- gredients of our food and also of the muscles and nerves of our bodies. (See Chapter XXIV.)

112. Preparation of nitrogen. — Nitrogen is prepared on a large scale from hquid air (129). It can also be ob-

NITROGEN — THE AIR

103

taincd from air by removing the oxygen from air by phosphorus (Fig. 56).

Phosphorus is put in a small dish or a cruci- ble cover supported on a cork floating in a vessel of water. Upon igniting the phos- phorus with a hot wire and placing a bell jar quickly over the cork, the phosphorus and oxy- gen unite, forming clouds of white phosphorus pentoxide (P2O5). This solid soon dissolves in the water, which rises inside the jar owing to the removal of the oxygen, and the nitro- gen is finally left.

Fig. 56. — Prepara- tion of nitrogen by burning phos- phorus in confined air

Nitrogen is prepared in the laboratory by heating a solution of sodium nitrite (NaNOo) and am- monium chloride (NH4CI). These two compounds form the unstable compound ammonium nitrite (XH4XO2) which decomposes into nitrogen (N2) and water (H2O) ; the gas is collected in bottles over water, just as in the case of oxygen and hydrogen. Small quantities of nitro- gen may be readily obtained by heating ammonium dichromate f(NH4)2Cro07) in a test tube.

113. Properties of nitrogen. — Nitrogen is a colorless gas, and has no taste or odor. It is a little lighter than oxygen and air. A liter at standard conditions weighs 1.25 gm. (A liter of oxygen weighs 1.43 gm. and one of air 1.29 gm.) It is only slightly soluble in water. Subjected to a low temperature and increased pressure, nitrogen be- comes a colorless liquid and ultimately a white solid.

114. Chemical conduct of nitrogen. — Nitrogen does not support combustion nor sustain Ufe. Flames are extin- guished by nitrogen and animals die in it, because the supply of ox>^gen is cut off. The fact that nitrogen quickly ex- tinguishes a candle flame and kills a mouse was first ob-

I04 PRACTICAL CHEMISTRY

served by Rutherford, a Scottish physician, who discovered the gas in 1772. Soon after, Lavoisier showed the true relation of nitrogen to the atmosphere ; and because the gas would not support life, he called the gas azote — a name now used by some French chemists. The name nitrogen was given to it because it is a constituent of the important compound niter {i.e. saltpeter, potassium nitrate, KNO3).

Nitrogen is very much less active chemically than oxygen. Indeed, if we test it, as we did oxygen and hydrogen, we find that it responds to none of the common tests. It is sometimes called an inert element, because it does not com- bine with elements at ordinary temperatures. At high temperatures and under special conditions, how^ever, ni- trogen forms many compounds. It combines with mag- nesium and a few other metals at red heat, forming ni- trides, e.g. magnesium nitride (Mg3N2). Electric sparks cause nitrogen to combine with oxygen and with hydro- gen, forming nitrogen oxides and ammonia (NH3). Both these reactions, if hastened by a catalyst, proceed rapidly enough to enable us to convert nitrogen from the air into compounds needed as fertihzers and explosives (196, 197) .

115. Uses of nitrogen. — Nitrogen on account of its inertness is used to fill some kinds of electric hght bulbs, and the stem of high-boiUng thermometers. It is also used in making ammonia, nitric acid, and a nitrogen fertilizer called calcium cyanamide (CaCN2).

116. Nitrogen and life. — Nitrogen, as well as oxygen, is vitally connected with life, though in a different way. All animals need nitrogen for their growth. But although we live in an atmosphere containing such a large proportion of this gas, we cannot assimilate it directly. The nitrogen we inhale (along with the oxygen) is exhaled again unused. The nitrogen needed by animals must be eaten in the

NITROGEN — THE AIR

105

form of nitrogenous food, such as lean meat, fish, wheat, and other grains. (See Protein, Chapter XXIV.)

Nor have plants, with few exceptions, power to assimi- late free nitrogen from the atmosphere. Most plants take up combined nitrogen from the soil in the form of nitrates or of ammonia. Hence combined nitrogen is being con- stantly removed from the soil. In order to restore it, some nitrogen compound must be added, e.g. sodium nitrate (NaNOs), calcium nitrate (Ca(N03)2), ammonium chloride (NH4CI), ammonium sulphate ((NH4)2S04), or calcium cyanamide (CaCXo) ; organic materials are often used, e.g. manure, dried blood, and meat or fish scraps. A substance or mixture which restores nitro- gen (or some other chemical element like phosphorus or potassium) to the soil is called a fertilizer.

Many experiments have shown, however, that legumi- nous plants, such as peas, beans, and clover, take up nitrogen from the air by means of bacteria, which are in nod- ules on their roots (Fig. 57). with a preparation which contains nitrogen-forming bacteria.

117. Air is a mixture and not a compound. — Air is a mixture of several gases. Oxygen, nitrogen, and argon are the three ingredients that are always present in nearly constant proportions. Variable proportions of water va- por and carbon dioxide gas are always found, and also small quantities of compounds related to ammonia and

Fig- 57- — A plant with nodules on the roots

Sometimes soil is treated

io6 PRACTICAL CHEMISTRY

nitric acid. Near cities the air may contain considerable dust, sulphur compounds, and acids ; at the ocean some salt is often found.

Hitherto we have dealt almost exclusively with com- pounds, and we have found that compounds have a constant composition, i.e. the constituents are united in a proportion which is always the same in the case of a given compound. Furthermore, we have seen that if we wish to make a com- pound, we must cause the elements to unite chemically, and conversely if we wish to decompose a compound we must tear the compound apart chemically. In a few words, chemical action is always concerned where compounds are formed or decomposed. Moreover, this action is usu- ally accompanied by heat changes. The following facts show that air is not a compound but a mixture of gases : —

(i) The proportion of oxygen and of nitrogen is not fixed, but varies between small hmits. Therefore air does not have a constant composition and cannot be represented by a formula.

(2) When nitrogen and oxygen are mixed in approxi- mately the proportions that form air, the product is iden- tical with air, but the act of mixing gives no evidence of chemical action.

(3) When air is dissolved in water, a larger proportion of oxygen than nitrogen dissolves. If the oxygen and nitrogen were combined, the dissolved air would contain the same proportions of oxygen and nitrogen as air itself.

(4) When air is liquefied and allowed to boil, the nitrogen boils off first (129). If air were a compound, liquid air would have a fixed boiling point.

118. Proportions of the constant ingredients of air. — For many years it was beheved that pure air consisted solely of oxygen and nitrogen. But in 1894 it was found

NITROGEN — THE AIR

107

that nearly 1.2 per cent (by volume) of the gas hitherto called nitrogen is argon (124). The normal proportions (by volume) of the constant ingredients of air are nitrogen 78.122 per cent, oxygen 20.941 per cent, and argon 0.937 per cent. These numbers are often stated approximately as oxygen 78, nitrogen 21, and argon 0.94.

119. " Composition " of air. — Although air is a mixture, we usually speak of its '' composition." Samples of air from various parts of the globe show such a remarkable uniformity in the propor- tions of the constant ingredients that chemists have fallen into the habit of applying the term composition to air.

The proportion of oxygen in the air can be found by several methods. In one, a known volum.e of air is shaken in a closed bottle with a mixture of pyrogallic acid and sodium hydroxide ; this so- lution absorbs the oxygen and leaves the nitrogen and argon unchanged. In another, a graduated glass tube, containing a known volume of air, is inverted in a jar of water, and a piece of phos- phorus attached to a wire is introduced into the tube (Fig. 58). The oxygen combines with the phos- phorus. In a few hours the phosphorus is removed, and the volume of residual gas is read. The dif- ference between the first and last volumes is oxy- gen. There is no simple way of separating the nitrogen and argon.

Fig. 58. — Find- ing the per cent of oxygen in air by phos- phorus

120. Water vapor in the air. — Water vapor is always present in air, owing to constant evaporation from the ocean and other bodies of water. When the temperature of the air falls sufficiently, the water vapor condenses and is de- posited in the form of dew, rain, fog, mist, frost, snow, sleet, or hail (67). The clouds are masses of minute drops of Uquid water formed by condensation of the water vapor

io8 PRACTICAL CHEMISTRY

in the cold upper air. The condensation of considerable moisture forms large drops, which fall to the earth as rain.

The total amount of water vapor in the lower regions of the atmosphere is very large. However, the proportion in the air of different regions varies between wide limits. In rainy regions, such as tropical countries, it is large, while in desert countries it is small.

A given volume of air absorbs a definite volume of water vapor and no more. Warm air holds more than cool air. Air containing its maximum amount of water vapor is said to be saturated at that temperature, or to contain loo per cent of water vapor. The saturation point is also called the dew point, i.e. the temperature at which it begins to deposit dew. On a pleasant day in a temperate climate the relative humidity, i.e. the relative amount of water vapor present, may vary from 30 to 90 per cent, the average being about 65 per cent.

121. Test for moisture in the air. — The presence of water vapor in the air is shown by the moisture which collects on the outside of a vessel containing cold water, such as a pitcher of iced water. The moisture comes from the air around the vessel. For a similar reason, water pipes in a cellar and the cellar walls themselves are moist in summer.

In the laboratory the presence of water vapor may be shown by exposing to the air a deliquescent substance, such as calcium chloride (86). If a known volume of air is drawn slowly through a weighed tube containing calcium chloride, the increase in weight gives the weight of water in the volume of air.

The relative humidity may be calculated by first ascertaining the dew point and then comparing the final (dew point) pressure and the original (observed) pressure. For example, if air at 20° C. must be cooled to 12° C. before it deposits dew, its moisture content would be expressed as 10.48 mm. (73). If it were saturated at 20° C, its moisture content would have been equivalent to 17.41 mm. That is, its relative humidity was 10.48 -h 17.41, or 60.19 per cent.

NITROGEN — THE AIR 109

122. Physical comfort depends on water vapor in the air. — The proportion of water vapor in the air has a marked effect on our bodily comfort. When the relative humidity is near the average, we feel comfortable. But if the air is moist or dry we are very uncomfortable.

The relation of water vapor to bodily comfort is simple. Our bodies have a normal and nearly constant temperature of 98.6° F. This temperature is maintained by the heat produced by the chemical changes in our bodies, espe- cially the oxidation of waste tissue by the ox\' gen carried by the blood to all parts of the body (27). But this tem- perature is regulated partly by radiation of heat and partly by evaporation of water from the surface of the body. Now if the air is moist, evaporation proceeds with difficulty ; whereas if the air is dry, it pro- ceeds too fast. In either case, we are uncomfortable, and try by various devices to increase our comfort. Thus, we use fans and w^ear thin clothing to promote evaporation, or we m.oisten the air by exposing pans of w^ater.

Our comfort and efficiency depend on maintaining con- stant, moderate evaporation from the surface of the body. This desirable condition may be accomplished, or at least favored, by living in the open air as much as possible, main- taining the average humidity in our houses, wearing cloth- ing adapted to cUmatic conditions, and ventilating the rooms in which we live.

123. Carbon dioxide in the air. — Carbon dioxide is one product of the breathing of animals, the combustion of fuels, and the decay of organic substances (33). By these processes vast quantities of carbon dioxide are being con- stantly introduced into the air. The proportion in ordinary- air is 3 to 4 parts in 10,000 parts of air, i.e. 0.03 to 0.04 per cent. In crowded rooms it is often as high as ^7, parts

PRACTICAL CHEMISTRY

in 10,000. The proportion of carbon dioxide in the at- mosphere as a whole is practically constant, owing to winds and air currents, and largely, also, to the fact that this gas is absorbed by all green plants (39).

The presence of carbon dioxide in air is detected by calcium hydrox- ide. If calcium hydroxide solution is exposed to air, the carbon diox- ide interacts with the calcium hydroxide, forming a thin, white crust of insoluble calcium carbonate on the surface of the Uquid. If con- siderable air is drawn through the calcium hydroxide solution, the liquid becomes milky, because the particles of calcium carbonate are suspended in the liquid (33 and Fig. 12).

124. Argon in the air. — Argon, as stated above (118), is an essential and constant ingredient of the air, the pro- portion being 0.937 P^^ cent by volume.

Argon is a colorless, odorless gas which is a httle heavier than oxygen. It dissolves in water to the extent of

about 4 volumes in 100. It can be hquefied and solidified.

A conspicuous property of argon is its lack of chemical activity. No compounds of this ele- ment have as yet been prepared or discovered. The name argon is hap- pily chosen, being derived from Greek words signi- fying inert.

125. Discovery of ar- gon. — Argon was de- tected and first studied Fig. 59. - Ramsay (1852-1916) in 1894 by Rayleigh and

NITROGEN — THE AIR iii

Ramsay (Fig. 59). Rayleigh found that nitrogen ex- tracted from air had a greater weight than an equal volume of nitrogen obtained from compounds of nitrogen. Consequently, they beheved that the nitrogen from air contained another gas hitherto overlooked. Experiments (by Ramsay) showed that after the oxygen and nitrogen were removed from purified air, there still remained a small quantity of a new gas. It was named argon and given the symbol A.

Argon was first obtained by Ramsay by passing pure air over heated copper to remove the oxygen, and then the remaining gas over heated magnesium or calcium to remove the nitrogen. He also passed elec- tric sparks through a mixture of air and oxygen, and absorbed the nitrogen oxides in potassium hydroxide solution. The latter method is a repetition of the one used by Cavendish in 1785 when he deter- mined the composition of air ; he observed and recorded the fact that a small bubble of gas always remained, which was doubtless argon. To Cavendish belongs the honor of first observing this element.

A recent method of obtaining argon consists in liquefying air, and allowing the nitrogen and then the argon to boil off from the ox>'gen. This method is based on the fact that liquid argon boils at a higher temperature (-186° C.) than nitrogen ( — 195.7° C).

126. Other gases in the atmosphere. — Helium, neon, krypton, and xenon are inert gases discovered by Ramsay subsequently to argon. With the exception of neon, they constitute an exceedingly minute proportion of the atmosphere. Like argon they do not form compounds. Ramsay estimated that in 1,000,000 parts of the atmos- phere there are 3 to 4 parts of helium, 10 to 20 of neon, 0.05 of krypton, and 0.006 of xenon.

127. Helium was detected in the atmosphere of the sun by Lockyer in 1868. It was found by Ramsay, soon after he discovered argon, in the gases expelled from cer- tain rare minerals and in the gas and water of some min- eral springs. Helium is now obtained from natural gas which issues from the ground in certain parts of the United

112 PRACTICAL CHEMISTRY

States. By liquefying the gas, the helium can be sepa- rated from the other ingredients, owing to its low boiling point (—269° C). It is a Ught, non-inflammable gas. Both of these properties make hehum an excellent gas for filling balloons and hghter-than-air airships, and success- ful experiments have been made along this line.

Helium is one of the disintegration products of radium (620) ,

128. What is liquid air? — Liquid air is a mixture of liquid oxygen and liquid nitrogen. It sometimes looks cloudy, owing to the presence of solid carbon dioxide and ice. When these soHds are removed by filtering, the fil- trate has a pale blue tint.

129. Properties of liquid air. — If a beaker is filled with liquid air, the latter boils vigorously, the surrounding air

becomes intensely cold, frost gathers on the beaker, and in a short time the liquid air will disappear into the air of the room. If, however, Hquid air is put into a Dewar flask, evaporation takes place so slowly that some liquid air will remain in the flask several days.

A Dewar flask (Fig. 60) consists of two flasks,

one within the other sealed together air-tight at

Fig. 60. — A ^Yie top ; the space between the flasks is a vacuum.

e\\ ar as ^j^^ surfaces of the flasks are coated with silver, which

reflects heat and helps retard the evaporation of the liquid air. Liquid

air is stored and transported in Dewar flasks.

Liquid air boils at about —190° C. If it is allowed to boil in a proper apparatus, the nitrogen (boiling point — 195.7° C.) escapes first, leaving more or less pure oxygen (boiling point —182.9° C). (See also 125.) The industrial separation is accomplished this way, the two gases being collected in separate tanks.

NITROGEN — THE AIR 113

130. Some experiments with liquid air. — A tin or iron

vessel which has been cooled by hquid air is so brittle that it may often be crushed with the fingers ; while a piece of rubber tubing becomes as brittle as glass. Mercury freezes so hard in liquid air that it can be used as a hammer to drive a nail.

When liquid air is poured into a kettle standing on a block of ice, the hquid air boils vigorously because the ice is so much ^' hotter." If a kettle of liquid air is placed over a hghted Bunsen burner, frost and ice collect on the bottom of the kettle, because the intense cold pro- duced by the evaporation of the liquid air in the kettle solidifies the water vapor and carbon dioxide, which are the tw^o main products of burning illuminating gas. If wa- ter is now poured into the kettle, the hquid air boils over and the water is instantly frozen ; the water is so much " hotter " than the liquid air that the latter boils more violently, and since its rapid evaporation causes the ab- sorption of heat, the water loses heat and becomes ice. ■ Ordinary liquid air is from one half to one fifth hquid oxygen, and will support combustion. A glowing stick or a red-hot rod of steel burns brilliantly in this cold hquid.

131. Manufacture of liquid air. — Liquid air is manufactured by forcing air cooled by water through a pipe to a valve. As it escapes through the valve, it expands and its temperature falls, because expansion is a cooling process. After expansion, the cold air is led back over the outer surface of the same pipe by which it came, where- upon it rapidly regains its former temperature. But in so doing it cools the pipe itself and the air within it. This latter air in turn expands and falls in temperature, but as it was colder than the first portion before expansion, so it is colder after expansion. Since the pressure within the pipe is maintained by a continuous supply of air under pressure, the pipe becomes continually colder, until finally the expanding air at the valve liquefies in part.

114 PR-\CTIC-\L CHEMISTRY

EXERCISES

1. How ii nitrogen prepared? :?ummarize its properties. Com- pare the chemical conduct of nitrogen and oxygen. For what is nitro- gen used?

2. What is the relation of nitrogen to the life of (a) animals and (6) plants?

3. What are the two chief ingredients of the atmosphere? The constant ingredients? The variable ingredients? The ingredients found in traces? What special substances are sometimes found in the air of cities ?

4. State the volumetric composition of air. How is it found?

5. Has air a chemical formula? If so, what is it? If not, why?

6. Describe the action of air upon (a) calcium hydroxide and

(b) calcium chloride.

7. What is the meaning of the term de-u- paint? Relaihe humidity ?

8. Discuss the relation of physical comfort to water vapor in air.

9. Give several proofs that air is a mixture.

10. What is argon? Give a brief account of (a) its discovery, (6) its prop>erties. (c) its method of preparation. What proportion of air is argon ? What is the significance of the name argon ?

11. What is liquid air? What are its chief properties? Describe a Dewar flask.

12. Topics for home study: (a) Helium, (h) Uses of nitrogen. ic) Function of oxj-gen in the air . (d) Atmospheric pressure, (ei Ven- tilation, (f) Carbon dioxide in the air. ig) Lavoisier and combustion. (A) Cavendish and air.

13. Practical topics: (a) How would you distinguish nitrogen from carbon dioxide ? (6) Why are mines and cellars damp ? (c) Why does the composition of the atmosphere var>- so slightly? (d) How can the relative humidity of a room be increased in a heated house ?

PROBLEMS

1. What is the weight of air in a room 6 X 8 X 5 m. ? !'A liter of air weighs 1.29 gm.)

2. How many kilograms of pure air are needed to yield [a) 100 kg. and (h) 100 1. of oxygen (at standard conditions) ?

3. Express in inches the following barometer readings : (a) 760 mm., (6) 745 mm., (c) 70 cm., (d) 0.769 m., (e) 7.49 dm., (/) 780 mm., (g) 5 mm.

4. What is the weight at 0° C. and 760 mm. of {a) 1000 cc. of dry air? Of {h) 95 L, (c) 95 cc, {d) 95 cu. m.?

CIL\PTER X CHEMICAL REACTIONS — EQUATIONS — CALCULATIONS

132. Chemical reactions. — The substances that par- ticipate in a chemical change are said to undergo chemical action, to interact, or to react. A single chemical change is called a chemical reaction, an interaction, or simply a re- action.

133. What is a chemical equation? — We learned in 14 that a reaction can be represented in a condensed form called an equation. Thus : —

Iron -f Sulphur = Iron Sulphide

In the preceding chapters certain chemical changes were represented by equations. In 105. 106 we found that ele- ments and compounds are represented by s\Tnbols and formulas. Therefore, we can now remodel the prelimi- nan' equations into chemical equations by using sjinbols and formulas in place of words. The above equation then

becomes : —

Fe — S = FeS

134. How to write an equation. — The mere change of words to symbols and fonnulas would not always give a correct equation. In order to write the equation that correctly expresses a chemical change, we must know cer- tain facts and express them in the correct chemical way. Thus, we must know the symbols and formulas of the factors (the reacting substances') and of the products (the hnal substances'). Next we must \sTite a preliminan- equation.

ii6 PRACTICAL CHEMISTRY

and then adjust it so that there is the same number of atoms of each element on both sides of the final equation.

These are the steps in writing an equation: (i) Write on the left (of the equality sign) the symbol or formula of each factor. (2) Write on the right the symbol or for- mula of each product. (3) Balance the preliminary equa- tion, if necessary.

To balance an equation we increase the number of atoms, or molecules, or both — as necessary — until the number of atoms of each element (free or combined) is the same on both sides of the equation. Let us take three examples.

First, when magnesium and oxygen form magnesium oxide, the preliminary equation is : —

Mg + O2 = MgO

Inspection shows that there are 2 atoms of oxygen (in O2) on the left but only i (in MgO) on the right. Hence we must increase the number of oxygen atoms on the right. We do this by prefixing the coefficient 2 to the formula MgO (not by altering the MgO). This change necessitates multiplying Mg on the left by 2. The final balanced equa- tion then becomes : —

2Mg + 02 = 2MgO

Final inspection shows that the same number of atoms of each element is on both sides of the equation, hence the equation is correct.

Second, when zinc and hydrochloric acid interact, hy- drogen and zinc chloride are formed. The preliminary equation is: —

Zn -f HCl = H2 + ZnCl2

By inspection it is evident that 2 atoms of chlorine (in ZnCl2) are on the right and only i is on the left. Hence we m.ust

CHEMICAL REACTIONS — EQUATIONS 1 1 7

multiply HCl by 2, thereby providing 2 atoms of chlorine, and giving also the 2 atoms of hydrogen. The final bal- anced equation becomes : —

Zn + 2HCI = H2 + ZnCl2

Third, when potassium chlorate is decomposed, oxygen and potassium chloride are produced. The preliminary equation is : —

KCIO3 = 02 + KCl

By inspection we see that there are 3 atoms of oxy^gen (in KCIO3) on the left but only 2 (in Oo) on the right. To obtain the same number on each side, we must multiply KCIO3 by 2 and O2 by 3. The second preliminary equation is : —

2KCIO3 = 3O2 + KCl

Inspecting again, it is clear that there are 2K and 2CI on the left but only i each on the right. Therefore, we balance by multiplying KCl by 2, and obtain the final cor- rect equation : —

2KCIO3 = 3O2 + 2KCI

135. Some precautions to be observed in writing equa- tions. — (a) Correct formulas must be used. If a formula is not remembered or is not given in the immediate text, it should be looked up in the book ; this can be done by finding the name of the substance in the index and con- sulting the proper page where, as a rule, the formula is given. Thus, when sodium interacts with water, sodium hydroxide and hydrogen are formed. By referring to the page in the text under which sodium hydroxide is indexed.

ii8 PRACTICAL CHEMISTRY

we find its formula is NaOH. The preliminary and final equations are : —

Na + HoO = NaOH + Ho 2Na + 2H2O = 2NaOH + H2

(b) It must not be overlooked that the correct formulas of many elementary ga5g5, such as oxygen, hydrogen, nitrogen, and chlorine, are O2, H2, N2, CI2 respectively (not O, H, etc.).

(c) In balancing an equation these rules are helpful : (i) Start with the formula containing the most atoms of one element. (2) Find the other formula (or symbol) containing this element, and increase the number of atoms of this element by prefixing a coefficient, not by altering the for- mula (or symbol) — see middle of page 116. (3) Increase again, if necessary. (4) Balance in the same way for the other elements. (5) Check up, so that finally the total num- ber of atoms of each element is the same on both sides. Let us take an example. When phosphorus burns in oxy- gen, phosphorus pentoxide is formed. The formula of phosphorus pentoxide is P2O5. Hence the prehminary equation is : —

P + O2 = P2O5 By inspection we see that P2O5 needs at least 5 atoms of oxygen. Clearly the only way to balance the equation for oxygen is to multiply O2 by 5 and P2O5 by 2 ; this ad- justment gives 10 atoms of oxygen on each side, thus : —

P -f 5O2 = 2P2O5 But multiplying P0O5 by 2 gives 4 atoms of P on the right, because a coefficient multipHes the whole formula. That is, 2P2O5 means two molecules each containing 2P and 5O. Hence we balance for P by multiplying P on the left by 4. The final equation then becomes : —

4P + 5O2 = 2P2O5

CHEMICAL REACTIONS — EQUATIONS 119

(d) Only the substances that actually take part in the chemical change should be included in the equation. Thus, when magnesium is burned in air, the nitrogen of the air does not unite with the magnesium (to any extent). Hence nitrogen does not appear in the equation. Similarly in the equation for the preparation of hydrogen from zinc and hydrochloric acid, no water (H2O) appears as a factor because the water (in the dilute sulphuric acid) does not participate in the reaction.

(e) The terms of an equation, i.e. symbols and for- mulas, cannot be transposed as in the case of an algebraic equation. A chemical equation is the expression of an actual chemical change ; it is the outcome of an experi- ment. A chemical equation, although it contains the plus sign, has none of the properties of an algebraic equation except equality of the total weights on each side.

136. How to read an equation. — The plus ( + ) sign may be read and or plus and the equality ( = ) sign fonn(s), gives{s), or equal{s). An arrow ( — >■) is sometimes used instead of the equality sign ( = ) ; both signs are read in the same way. Since equations are really expressions of equal- ity between two total weights, there is a good reason for using the sign of equality. (Compare 174.)

Consider the equation : —

Zn + H2SO4 = H.2 + ZnS04

This equation may be read in several ways: (i) zinc and sulphuric acid form (or give) hydrogen and zinc sulphate ; (2) zinc plus sulphuric acid equal hydrogen plus zinc sul- phate ; (3) one atom of zinc and one molecule of sulphuric acid form one molecule of hydrogen and one molecule of zinc sulphate. (A fourth way is given in 208.)

I20 PRACTICAL CHEMISTRY

137. Ordinary equations have a limited meaning. — The equa- tions we are studying in this chapter might be called ordinary equa- tions or atomic equations. They merely show the result of a reaction by the proper number of atoms and molecules. Other kinds of equa- tions are used to tell other facts about chemical change, and these will be discussed in the proper places.

Ordinary equations show nothing about the physical conditions under which a reaction occurs. Thus, the equation

2Mg +02= 2MgO

does not show that the reaction takes place quickly at a moderate temperature {e.g. heat from a burning match). Nor does the equation

2HgO = 2Hg + O2

show that the reaction takes place rather slowly and only by con- stant heating at a high temperature {e.g. heat from a Bunsen burner). So also ordinary equations do not include the solvent, although many reactions occur only in solution. Moreover, ordinary equa- tions tell nothing about the way in which a reaction occurs. Thus, in the equation

2KCIO3 = 3O2 + 2KCI

several chemical changes doubtless occur which do not appear in the equation. The main purpose of this equation is to express in the simplest way the transformation of potassium chlorate into the two products oxygen and potassium chloride.

138. Ordinary equations are gravimetric equations.—

Ordinary equations, in spite of certain limitations, have a practical use which is very important.

We have seen so far that equations show in a compact way the products formed by the reaction of certain sub- stances and that they also show the number of mole- cules (and in some cases the atoms) of each substance in- volved in a reaction. Thus, the equation

4H2O + 3Fe = Fe304 + 4H2

CHEMICAL REACTIONS — EQUATIONS 121

not only shows that iron oxide and hydrogen are the prod- ucts of the interaction of water (in the form of steam) and iron, but also that 4 molecules of water and 3 atoms of iron interact to form i molecule of iron oxide and 4 molecules of hydrogen.

Equations are quantitative expressions. They are based on the fact that a chemical change conforms to the law of the conservation of matter, viz., No matter is lost or gamed in a chemical change. We start with a certain total weight and we have the same total weight at the end of the re- action. This is only another way of saying that we have the same number of atoms on each side of a balanced equa- tion, because we have the same sum of atomic weights on each side of the equation. Indeed, there could be no " equation " unless both sides were equal in weight. Briefly, ordinary equations are gravimetric.

139. Calculations based on equations. — Each atom stands for a certain weight. Hence we can wTite the equa- tion for the interaction of zinc and sulphuric acid thus : —

Zn + H2SO4 = H2 + ZnS04

65 2 + 32+64 2 65+32+64

65 98 2 161

The equation in this form is read : 65 parts by weight of zinc plus 98 parts by weight of sulphuric acid equal 2 parts by weight of hydrogen and 161 parts by weight of zinc sulphate. That is, these numbers are the relative weights of the different substances involved in this reaction. Zinc and sulphuric acid always interact in the ratio of 65 to 98 and produce hydrogen and zinc sulphate in the ratio of 2 to 161.

In actual practice, of course, we do not have to start with exactly 65 gm. of zinc or with 98 gm. of sulphuric acid.

122 PRACTICAL CHEMISTRY

We can use any convenient \yeights ; but whatever weights we use, these two substances interact in the ratio of 65 to 98 and the quantity of either substance greater than the amount for the required ratio will be left unchanged.

Suppose we start with 45 gm. of zinc, pour sulphuric acid upon it, and let the reaction continue until the zinc is used up. We can calculate readily the weight of the sulphuric acid used.

First, we write the equation. Thus : —

Zn + H0SO4 = Ho + ZnS04

Second, we place under each term of the equation its correct relative weight, i.e. the weight for which the com- plete symbol or formula stands, using for this purpose the approximate atomic weights in the table on the back in- side cover. Thus : —

Zn + H2SO4 = Ho -f ZnS04

65 2 + 32 -f- 64 2 65 -f- 32 + 64 65 98 2 161

Third, we place above the symbol for zinc the given weight, which is 45 in this example, and above the for- mula for sulphuric acid the letter x to denote the required weight. Thus : —

45 ^

Zn -f- H0SO4 = Ho -t- ZnS04

65 98 2 161

Fourth, we state in a proportion the four quantities in- volved. In this proportion the equation weights (below the equation) are the first and second terms, while the corresponding known and required weights (above the equation) are the other two terms. Thus : —

65 : 98 : : 45 : a;

CHEMICAL REACTIONS — EQUATIONS 123

Fifth, we iinally solve the proportion for x, remembering that the product of the means (the two inner terms) equals the product of the extremes (the two outer terms) . Thus : —

98 X 45 = 65 X A- .-. X = ^}^AS^ 01- 5^ g ^^^^ 5^ 3 „^

The quantity of each product can be found by a similar proportion.

Another example will make the process clearer. This time we start with 15 gm. of potassium chlorate and wish to know the weight of oxygen which can be obtained. Pro- ceeding as above, we have these steps : —

(i) Chemical equation: —

2KCIO3 = 3O2 + 2KCI

(2) Relative weight equation : —

2KCIO3 = 3O2 + 2KCI

2(39 + 35-5+48) 3(32) 2(39 + 35.5) 2(122.5) 96 2(74.5)

245 96 149

(3) Reacting weight equation : —

15 X

2KCIO3 = 3O2 + 2KCI 245 96 149

(4) Proportion : —

245:96:: 15:0;

(5) Solution : —

96 X 15 = 245 X X :. X = 96_X_15^ ^^ ^ g^^ ^^^^ ^ g^ gni^

245 140. Calculations involving weight and volume. — Cal- culations like those in 139 are limited to weights — gravi-

124 PRACTICAL CHEMISTRY

metric calculations. But many reactions involve gases, and it is often necessary to know the volume of the gas as well as its weight. To calculate the volume, we calculate the weight in the usual way and then convert the weight into volume ; or vice versa.

Let us take two examples. First, suppose we wish to know the volume of hydrogen liberated by the interaction of 20 gm. of zinc and sufficient dilute sulphuric acid. Pro- ceeding as in 139, we have : —

(l)

Zn + H2SO4

= Ho + ZnS04

(2)

Zn + H2SO4 65 98

= H2 + ZnS04 2 161

(3)

Zn + H2SO4

65 98

= H2 + ZnS04 2 161

(4)

65:2::

20: X

(5)

2 X

20 = 65 X X :

•-^^/%

, or 0.61 S-

Then

(6) Since i liter of hydrogen weighs 0.09 gm. (52), the volume of hydrogen is : —

^^^= 6.83. Ans. 6.83 1. 0.09

As a second example, suppose we wish to know the weight of calcium carbonate needed to produce 200 liters of car- bon dioxide. First, we find the weight of 200 Hters of car- bon dioxide. Since i hter weighs 1.98 gm. (36 and Ap- pendix, § 3), the weight of 200 hters is 200 X 1.98, or 396 grams. Proceeding as above we have : —

(i) CaCOa + 2HCI = CO2 + CaCl2 + H2O

CHEMICAL REACTIONS — EQUATIONS 125

(2) CaCOs -f 2HCI = CO2 + CaClo + HoO 40+12+48 2(1+35.5) 12+32 40 + 71 2 + 16

100

(3) ^ CaCOa 100

44 396

III

+ 2HCI = CO.2 + CaClo + HiO

(4) (5)

100 : 44

44 X : 396

III

44 X X = 100 X 396 .-. X = ^^, or goo.

Ans. 900 grams.

141. Equations for preceding reactions. — The equations cor- responding to many reactions already discussed may be collected here, partly for review and partly for use in solving problems.

2HgO Mercuric Oxide

O2 Oxygen

2Hg

Mercury

2KCIO3 Potassium Chlo-ate

3O2 Oxygen

+ 2KCI Potassium Chloride

2Pb02

Lead Dioxide

O2 Oxygen

2PbO Lead Oxide

2Ba02 Barium Dioxide

Oxygen

2BaO Barium Oxide

2H.2O Water

2 Ho Hydrogen

0,> Oxygen

Sulphur

O2 Oxygen

SO2 Sulphur Dioxide

Carbon

O2 Oxygen

CO2 Carbon Dioxide

2CU

Copper

O2 Oxygen

2CuO Copper Oxide

Zn Zinc

H2SO4 Sulphuric Acid

II.. Hydrogen

ZnS04

Zinc Sulphate

126

PRACTICAL CHEMISTRY

Zinc

+ 2HCI

Hydrochloric Acid

H2 + ZnCls

Hydrogen Zinc Chloride

2Na +

Sodium

2H2O Water

H2 Hydrogen

+ 2XaOH Sodium Hydroxide

Calcium

2H2O Water

Hydrogen

+ Ca(0H)2 Hydroxide Calcium

2AI + 6XaOH = Ho

Aluminium Sodium Hydroxide Hydrogen

+ 2Xa3A103 Sodium Aluminate

2H2 + O2

Hydrogen Oxygen

2H2O

Water

CuO + Ho

Copper Oxide Hydrogen

NaCl + AgNOs

Sodium Silver

Chloride Nitrate

H2O Water

AgCl

Silver

Chloride

Copper

+ NaNOs Sodium Nitrate

EXERCISES

1. Prepare a summary of 132 to 138.

2. Define and illustrate the terms reaction and equation.

3. State (a) the three steps in writing an equation, and {h) the rules for balancing an equation.

4. What precautions must be observed in writing equations?

5. Select an equation from 141 and read it in three different ways.

6. Interpret the equation 2Cu + Oo = 2CuO by stating {a) what it means and {h) what it does not include or express.

7. Write equations for the following reactions : (a) Calcium and hydrochloric acid form calcium chloride and hydrogen, {h) Potas- sium sulphate and barium chloride form barium sulphate and potassium chloride, (c) Calcium carbonate and hydrochloric acid form calcium chloride, water, and carbon dioxide.

8. As in Exercise 7 : (a) Calcium oxide and carbon dioxide form calcium carbonate, {h) Chlorine and aluminium form aluminium tri- chloride, (c) Carbon and lead oxide (PbO) form lead and carbon mon- oxide.

CHEMICAL REACTIONS — EQUATIONS 127

9. What is a gravimetric equation? Illustrate.

10. Balance these equations : (a) BaCl2 + H2SO4 = BaS04 + HCl; (h) Pb(N03)2 + H2S = PbS + HNO3 ; (c) AICI3 + NH4OH = A1(0H)3 + NH4CI; (d) NaOH + CO2 = Na2C03 + H2O.

11. Balance these equations: (a) Zn + HNO3 = H2 + Zn(X03)2; (b) HCl + ZnO = ZnCl2 + H2O ; (c) H2SO4 + NaNOa = Na2S04 + HNO3; (d) SO2 + O2 = SO3.

12. State in order the steps in calculating the weight of a substance when the weight of only one reacting substance is given (139).

PROBLEMS

1. How many grams of oxygen can be prepared from (a) 45 gm. of mercuric oxide, {b) i kg. of potassium chlorate, (<:) 1000 gm. of water?

2. As in Problem 1, from (a) 750 gm. of lead oxide (PbOa), (b) 2200 gm. of barium dioxide (Ba02)?

3. Hydrogen is prepared from sulphuric acid and 40 gm. of zinc. Calculate the weights of the products of the reaction.

4. If a balloon holds 150 kg. of hydrogen, how much (a) zinc and

(b) sulphuric acid are needed to generate the gas?

5. What volum.e of oxygen at standard conditions could be ob- tained from 10 gm. of potassium chlorate?

6. If 10 gm. of pure carbon are burned in air, what weights of other substances are involved?

7. One gram of copper is heated intensely in air, and the product is reduced by a gas. Calculate (a) the weights of the other substances involved in the two reactions, and (b) the volume of the gas in the second reaction.

8. How many grams of potassium chlorate are needed to prepare (a) 100 gm. of oxygen and (b) 100 1. (at standard conditions)?

9. Calculate the weights needed in the following reactions : (a) water and 100 milligrams of sodium, (b) calcium and 100 milligrams of water, (c) sodium hydroxide and 25 gm. of aluminium.

10. What weight of carbon dioxide is formed by burning a ton of coal which is 90 per cent carbon?

11. Suppose 85 gm. of water are decomposed. What (a) weights and (6) volumes of gases are produced?

12. Sixty grams of mercuric oxide are decomposed. What volume of oxygen at 91° C. and 380 mm. is produced?

13. How much water can be obtained from (a) 34 gm. of crystallized zinc sulphate (ZnS04.7H20), (b) 1000 kg. of selenite (CaS04.2H20),

128 PRACTICAL CHEMISTRY

14. Write the equation for the interaction of barium nitrate and sodium sulphate. If 170 gm. of barium nitrate are used, calculate the weights of the other compounds involved.

15. Ammonia and hydrogen chloride form solid ammonium chloride. Write the equation for this reaction. If 210 gm. of ammonia are used, calculate (a) the weights of the other compounds involved, and (b) the volumes of ammonia and hydrogen chloride needed.

16. The oxygen is liberated from 10 gm. of potassium chlorate, and ID gm. of sulphur are burned in the gas. How much sulphur, if any, is left?

CHAPTER XI

CHLORINE - HYDROGEN CHLORIDE— HYDROCHLORIC ACID

.^^fl

142. Introduction. — So far we have studied the ele- ments oxygen, carbon, hydrogen, and nitrogen, the com- pounds carbon dioxide, carbon monoxide, and water, and the mixture air. In this chapter we shall study the element chlorine and its common compound hydrogen chloride.

143. Occurrence of chlorine. — Free chlorine is never found in nature, but its compounds, especially chlorides, are widely distributed, the most abundant being sodium chloride (NaCl) or common salt. Many compounds of chlorine with potassium, magnesium, and calcium are found in the deposits at Stassfurt in Germany. (See Potas- sium, Chapter XXIX.) Sea water contains sodium and magnesium chlo- rides.

144. Preparation of chlorine. — Chlorine is prepared on a large scale by the electrolysis of a solution of sodium chloride. When an electric current is passed through a solution of sodium, chloride, chlorine gas is liberated in one compartment of the apparatus and sodium hydroxide is formed in the other. The chlorine is conducted off through pipes, and the dis- solved sodium hydroxide is drawn off at intervals. This process is further described in 456.

129

Fig. 61. — Apparatus to illustrate the prep- aration of chlorine by the electrolysis of a solution of sodium chloride.

I30

PRACTICAL CHEMISTRY

This process can be readily demonstrated. The apparatus is shown in Fig. 6r. A solution of sodium chloride is put in the battery jar .4 ; a little litmus solution is added and then enough dilute hydro- chloric acid to color the solution a distinct red. A block {B) divides the jar into two compartments (C and D), and the two pieces of elec- tric light carbon serve as electrodes {E and F). Soon after the cur- rent (from four or more cells or from a reduced street current) is turned on, the solution is bleached by the liberated chlorine in one compartment and turned blue by the sodium hydroxide in the other. The chlorine can also be detected by its odor.

Chlorine is prepared in the laboratory by heating con- centrated hydrochloric acid with manganese dioxide. The equation is : —

4HCI + Mn02 = CI2 + MnClo + 2H2O

Hydrochloric Acid

Manganese Dioxide

Chlorine

Manganese Bichloride

Water

A simple form of laboratory appara- tus is shown in Fig. 62. The manga- nese dioxide is put into the flask A and concentrated hydrochloric acid is introduced through the dropping tube B. By gently heating the flask, chlo- rine passes to the bottom of the bottle G and displaces the air.

145. Properties of chlorine. —

Chlorine is a greenish yellow gas. Its color suggested the name chlorine (from the Greek word chloros, meaning greenish yellow), which was given to it by paring chlorine in the lab- Davy about 1810. It has a dis- oratory agreeable odor, which is very pen-

etrating. If breathed, it irritates the sensitive lining of the nose and throat ; a large quantity produces lung trouble, and would ultimately cause death. Chlorine was the first

Fig. 62. — Apparatus for pre-

CHLORINE — HYDROGEN CHLORIDE

131

poison gas used in the World War. It is about 2.5 times as heavy as air. Hence it is easily collected by downward displacement, i.e. by conducting it downward to the bot- tom of a vessel and allowing it to displace the air. A liter of dry chlorine at 0° C. and 760 mm. weighs 3.22 gm.

Fig. 63. — Interior of a plant for liquefying and sioring chiorme. Cylinders full of chlorine are in the foreground

Chlorine can be readily liquefied and solidified. Liquid chlorine is sold in strong iron cylinders (Fig. 63 and ?lso

Fig. 3S)-

146. Chlorine dissolves in water. — The solution of chlorine is yellowish, and smells strongly of chlorine. Chlorine water, as the solution is called, is unstable. If the solution is placed in the sunhght, oxygen is slowly liberated and can be collected in a suitable apparatus, e.g. a long tube (87 and Fig. 46).

The oxygen does not come directly from the water but from a compound called hypochlorous acid (HCIO), which is formed in small quantities in the solution ; some hydrochloric acid is also formed. The reaction may be represented thus : —

132 PRACTICAL CHEMISTRY

CI2 -f H2O = HCIO + HCl

Chlorine Water Hypochlorous Acid Hydrochloric Acid

The hypochlorous acid is unstable and decomposes, thus : —

2HCIO = Oo + 2HCI

The reactions continue until all the chlorine is used up ; the equation for the completed change is : —

2HOO

2CU

4HCI +

Water

Chlorine

Hydrochloric Acid

Ox>'gen

147. Chemical conduct of chlorine. — Chlorine is a very active element. It unites vigorously with many elements at ordinary temperatures. Thus, if sodium, iron, copper, or other metals are merely warmed and then put into chlorine, they unite with the chlorine at once ; the sodium produces a dazzling light, and the copper and iron glow and emit dense fumes. These chemical changes illustrate the broad use of the term combustion. (See 53.) The compound formed in each case is a chloride, i.e. a com- pound of chlorine and one other element, e.g. sodium chlo- ride (NaCl) and iron chloride, (FeCls).

Chlorine combines readily with hydrogen (53). The equation for this reaction is : —

H2 + CI2 = 2HCI

Hydrogen Chlorine Hydrogen Chloride

A jet of burning hydrogen when lowered into chlorine continues to burn, forming a colorless gas called hydrogen chloride ; the latter becomes a white cloud when the breath is blown gently across the mouth of the vessel (155).

The tendency of chlorine to combine with hydrogen is so great that the hydrogen of many compounds is withdrawn chemically by chlorine. Thus, when cotton saturated with hot turpentine (CioHig) is put into chlorine, white fumes, due to the formation of hydrogen

CHLORINE — HYDROGEN CHLORIDE 133

chloride, appear almost at once ; soon the chemical change is so vig- orous that a flame is produced. Since carbon does not unite directly with chlorine, the white fumes of hydrogen chloride are finally ob- scured by a dense cloud of black smoke.

148. Chlorine bleaches. - Chlorine changes many col- ored substances into colorless ones. For example, colored cloth, whether dyed or naturally colored by impure sub- stances, is whitened {i.e. bleached) by moist chlorine. Dry chlorine does not bleach.

149. Sources of chlorine for bleaching. — Chlorine water, prepared by bubbling chlorine through cold water, is sometimes used as a bleaching agent. The commonest source, however, is bleaching powder (CaOClo) or, as it is often called, chloride of lime. It is a yellowish white substance which smells like chlorine, but the smell is really due to hypochlorous acid.

Bleaching powxler is manufactured by treating slaked lime (Ca(OH)o) with chlorine, the equation for the reaction being : —

Ca(OH)o +

Clo

CaOCl, +

H,0

Calcium Hydroxide

Chlorine

Bleaching Powder

Water

If bleaching powder is treated wdth an acid, chlorine is liberated, thus : —

CaOCI,

+ H0SO4 =

= Clo -f CaS04 +

HoO

Bleaching

Sulphuric

Chlorine Calcium

Water

Powder

Acid

Sulphate

These reactions also take place : —

CU + HoO = HCIO -f HCl HCIO = HCl + O

The bleaching action is really due to the liberated oxygen, which in this equation is represented by O, rather than O2. The formula O2 represents oxygen gas — two atoms

1.^4 PRACTICAL CHEMISTRY

united into one molecule. Whereas O represents an un- comhined atom of oxygen ; the free atom is in a more active chemical state, called the nascent state, because an atom of oxygen just liberated from a compound is ready, so to speak, to oxidize.

Another source of chlorine for bleaching is sodium hypo- chlorite (NaClO). This compound is prepared by passing a current of electricity through sodium chloride solution and allowing the products to react, thus : —

2NaOH

+ ci,

= NaClO + NaCl + HoO

Sodium

Chlorine

Sodium Sodium Water

Hydroxide

Hypochlorite Chloride

Sodium hypochlorite is unstable, and a cold dilute solution contains hypochlorous acid, owing to the interaction of the compound with water, thus : —

NaClO +

HoO

HCIO +

NaOH

Sodium

Water

Hypochlorous

Sodium

Hypochlorite

Acid

Hydroxid(

The hypochlorous acid furnishes nascent oxygen. Sodium hypochlorite solution is used by up-to-date laundries.

150. The process of bleaching. — Bleaching is an ox- idizing process. Oxygen, as we have just seen, is liberated in the nascent state from hypochlorous acid and readily decomposes the colored substances and changes them into colorless compounds.

A conventional diagram of the process of bleaching cotton cloth is shown in Y\g. 64. The pieces arc sewed together end to end in long strips and drawn by machinery from the roll .1 successively through vats containing bleaching powder solution B, weak acid C, and water D. At some point toward the end of the process the cloth passes through a vat containing acid sodium sulphite solution (or

CHLORINE — HYDROGEN CHLORIDE

135

similar mixture) E, called the antichlor, to remove traces of hypo- chlorous acid. After thorough washing, the cloth is dried and ironed by passing over hot cylinders FG , and finally wound on the roll H.

The chemical process of bleaching can be shown in the laboratory by a simple experiment. Put a little bleaching powder in the beaker

Fig. 64. — Diagram of the process of bleaching cotton cloth

A (Fig. 65), fill the beaker about one third full of water, and mix well. Fill the beaker B one third full of dilute sulphuric acid ; and the beaker C one third full of water. Press the lower half of the strip of colored cloth into the bleaching powder solution and then into the acid, passing it back and forth several times. Finally wash the cloth thoroughly in the water in C. The bleached cloth should have the general appearance shown in Fig. 65.

151. Uses of chlorine. —

Besides the use of chlorine in the manufacture of bleach- ing powder and

bleaching

Fig. 65. — Bleaching colored cloth

Fig. 66. — Fire ex- tinguisher con- taining carbon tetrachloride

mixtures, large Ciuantities of the gas are made into use- ful compounds of chlorine. One is carbon tetrachloride (CCI4), which is used in ''pyrene" fire ex- tinguishers (Fig. 66), and also as a solvent for extracting greases ; the non-inflammable cleaning mixture called '' carbona " con- tains carbon tetrachloride. Chlorine is also used in making disinfectants and

purifying water (69 and Fig. 35).

136

PRACTICAL CHEMISTRY

152. Hydrogen chloride and hydrochloric acid. — Hydrogen chloride is a gas, which is very soluble in water. Hydrochloric acid is the common name of a water solution of hydrogen chloride. This solution is known commer- cially as muriatic acid (from the Latin word muria, meaning brine), but it is more properly called hydrochloric acid. Hydrogen chloride is often called hydrochloric acid gas.

153. Preparation of hydrogen chloride. — This gas is prepared from sodium chloride by heating with sulphuric acid. If the mixture is gently heated, the chemical change is represented thus : —

NaCl

Sodium Chloride

H2SO4 =

Sulphuric Acid

HCl

Hydrogen Chloride

HXaS04

Acid Sodium Sulphate

But at a high temperature the equation is 2NaCl + H2SO4 = 2HCI

+ Na2S04

Sodium Sulphate

The solution is prepared by passing the gas into water.

154. Hydrochloric acid is manufactured in enormous quantities by a method essentially like that used in the laboratory.

The mixture of salt and sulphuric acid is put into the cast iron retort .4 (Fig. 67) and heated by the furnace 5 to a moderate temper- ature ; as soon as the >^' mass becomes pasty it

is raked out upon the flat heater A' and heated to a high tem- perature by the fur- nace B'. The hydro- gen chloride escapes

Fig. 67. — Sketch of the apparatus for the through C and C into manufacture of hydrochloric acid an absorbing tower

CHLORINE — HYDROGEN CHLORIDE 137

filled with resistant material over which water trickles ; as the gas passes up the tower, it is absorbed by the descending water, and the solution flows out at the bottom as concentrated acid.

Hydrochloric acid is also a by-product in the manufacture of sodium carbonate by the Leblanc {process (450).

155. Properties of hydrogen chloride. - Hydrogen chlo- ride is a colorless gas. It has a choking, sharp taste, and irritates the Hning of the nose and throat. It is about 1.25 times as heavy as air. A hter at 0° C. and 760 mm. weighs 1.64 gm. The gas becomes a colorless liquid when subjected to pressure and a moderately low temperature. The marked solubility of hydrogen chloride in water is one of its most striking properties. Even when it escapes into moist air, it forms white fumes which are really minute drops of a solution of the gas in the moisture of the air (147). At ordinary temperatures about 500 1. of gas dissolve in I 1. of water.

The solubility of hydrogen chloride in water can be shown by a simple experiment. The apparatus is arranged as in Fig. 68. The flask A is filled with hydrogen chloride. The medicine dropper B is partly filled with water, the stopper with its tubes is inserted, and the flask is then arranged as shown in the figure. By pinching the bulb of the dropper, a few drops of water are forced into the flask. This small quantity of water dissolves so much gas that a partial vacuum is formed in the flask ; pressure within the flask is reduced so much that the atmospheric pressure forces water from the jar C up the tube D and through the small opening into the flask.

156. Chemical conduct of hydrogen chloride. — Hy- drogen chloride does not burn nor support combustion. It is a very stable compound and can be heated to about 1800° C. before it begins to decompose. Perfectly dry hydrogen chloride has httle or no chemical activity. Whereas the moist gas unites readily with certain sub- stances, e.g. ammonia gas ; in this case dense white clouds of

138

PRACTICAL CHEMISTRY

0=<s

ammonium chloride (NH4CI) arc formed. This reaction is sometimes used as a test for hydrogen chloride (173).

157. Composition of hydrogen chloride. — Experiments show that hydrogen chloride is composed of hydrogen and chlorine in the ratio of I to I by volume. When a mixture of equal volumes of hydrogen and chlorine is ex- posed to the direct sunlight or to the action of an electric spark, the gases combine, hy- drogen chloride is formed with no residue, and the volume of the resulting gas equals the sum of the volumes of hydrogen and chlo- rine used. We express the volumetric rela- tions of hydrogen, chlorine, and hydrogen chloride by sa>mg: i volume of hydro- gen + I volume of chlorine = 2 volumes of hydrogen chloride. In the form of an equation, this fact becomes : —

Fig. 68. — Hydrogen chloride fountain

H2 + Clo = 2HCI

I vol. of I vol. 2 vols, of

Hydrogen Chlorine Hydrogen Chloride

The simple volumetric relation of these three gases is a good illustra- tion of Gay-Lussac's law (93, 176).

158. Hydrochloric acid. — The water solution of hy- drogen chloride is the common substance hydrochloric acid. The ordinary kind contains from 30 to 40 per cent (by weight) of hydrogen chloride, and is called concen- trated hydrochloric acid. Its specific gravity is about 1.2. In the laboratory we also use dilute hydrochloric acid, which is prepared from the concentrated acid by adding 4 to 5 times its volume of water.

Hydrochloric acid, like most members of the important class of compounds called acids, has a sour taste and reddens

CHLORINE — HYDROGEN CHLORIDE 139

blue litmus. It reacts with many metals, liberating hy- drogen and forming chlorides of the metals, thus : — Zn + 2HCI - Ho + ZnCl2

Zinc Hydrochloric Hydrogen Zinc

Acid Chloride

It also forms chlorides by interaction with oxides and hydroxides of metals, thus : —

CaO + 2HCI Calcium Oxide

= CaClo +

H,0

Calcium

Water

Chloride

= NaCl +

H2O

Sodium Chloride

Water

CaCOa =

-- CaCh

+ CO2 + H2O

Calcium

Carbon Water

Carbonate

Chloride

Dioxide

NaOH + HCl

Sodium Hydroxide

It interacts readily with carbonates. The equation in the case of calcium carbonate is : —

2HCI +

Hydrochloric Acid

Hydrochloric acid is an indispen- sable compound, and is used in many industrial processes. Like other acids, it is sold in bottles holding 2.5 hters and in securely packed glass car- boys containing 10 or more gallons (Fig. 69).

159. Aqua regia. - Hydrochloric ^g. ,^. _ ^ c^arboy acid and nitric acid interact and lib- of hydrochloric acid erate chlorine, thus : — packed for shipment

MURIATIC ^

f;ii) J

3HC1 +

HNO3

= 2CI + NOCl

2H2O

Hydrochloric

Nitric

Nascent Nitrosyl

Water

Acid

Chlorine Chloride

A mixture of one volume of concentrated nitric acid and three volumes of concentrated hydrochloric acid is usually

I40 PRACTICAL CHEMISTRY

used. If such a mixture is added to a metal, the nascent chlorine (2 CD forms a chloride of the metal. (Compare with nascent oxygen, end of 149.) The alchemists named the mixtures aqua regia, meaning " royal water," to em- phasize the fact that it dissolves the '' noble " metal gold. Another name is nitro-hydrochloric acid.

160. Chlorides are compounds of chlorine with other elements. They are formed, as we have already seen, by the direct combination of chlorine and metals (147) and by the interaction of hydrochloric acid with metaUic oxides or hydroxides (158). Most chlorides are soluble in water. But the chlorides of lead (PbClo), silver (AgCl), and one of the chlorides of mercury (HgCl) are not soluble; they are formed as insoluble solids, when hydrochloric acid or a soluble chloride is added to a solution of a lead com- pound, silver compound, or the proper mercury compound. Thus : —

Pb(N03)2 + 2HCI = PbCL -i- 2HNO3

Lead Nitrate Hydrochloric xVcid Lead Chloride Nitric Acid

The formation of insoluble solids by double decompo- sition (and certain other changes) is called precipitation, and the soUd itself is called a precipitate. Precipitates often have properties which are readily determined. Thus, silver chloride is white and curdy, and soon turns purple in the light ; moreover it dissolves in ammonium hydrox- ide owdng to the formation of a complex soluble compound, which, however, is transformed by dilute nitric acid into silver chloride. Other chlorides have difTerent properties. Hence, the precipitation of silver chloride serves as a test for hydrochloric acid and soluble chlorides.

A molecule of a chloride may contain one or more atoms of chlorine and occasionally the name of the compound indicates this fact, e.g.

CHLORINE— HYDROGEN CHLORIDE 141

manganese dichloride (MnCl-j), antimony trichloride (SbCls), carbon tetrachloride (CCU). If the same metal forms two chlorides, the two are distinguished by modifying the name of the metal ; the one containing the smaller proportion of chlorine ends in -ous, that con- taining the larger in -ic. Thus, mcrcurous chloride is HgCl, and mercuric chloride is HgClo.

EXERCISES

1. Prepare a summary of (a) chlorine and (b) hydrochloric acid.

2. Sketch from memory the apparatus used to prepare chlorine

(a) on a large scale, and (b) in the laboratory.

3. How can chlorine be quickly distinguished from the gases pre- viously studied ?

4. Summarize the chemical conduct of chlorine.

6. Topics for home study, (a) Nascent state, (b) Chlorine water, (c) Liquid chlorine, (d) Chlorine is an oxidizing agent.

6. What is (a) muriatic acid, (b) chloride of lime, (c) bleaching powder?

7. State the equation for (a) preparation of chlorine, (b) interaction of chlorine and water, {c) decomposition of hypochlorous acid, (d) prep- aration of hydrogen chloride, (e) interaction of hydrogen chloride and ammonia, (/) interaction of sodium chloride and silver nitrate.

8. Write an essay on " bleaching with chlorine."

9. Complete and balance these equations : (a) XaCl + H2SO4

= + Xa2S04 ; (b) H2 + = HCl ; (c) HCl + Xa.COs =

+ CC2 + ; (d) HgXOs H = HgCl + HXO3.

PROBLEMS

1. Calculate the weight of chlorine in (a) 2 kg. of sodium chloride.

(b) 2 mg. of calcium chloride, (f) i gm. of aluminium chloride.

2. How many grams of each product are formed when hydrochloric acid interacts with 85 gm. of manganese dioxide?

3. (a) What is the weight of 15 1. of chlorine gas measured at 20° C. and 790 mm. ? (b) How many grams of potassium chloride are needed to prepare the weight of chlorine found in (a) ?

4. How many grams of hydrogen chloride can be obtained from 27 gm. cf sodium chloride? How many liters (standard conditions)?

5. Calculate the percentage composition of (a) KCl, (b) CaCU. '

6. How much sodium chloride can be formed by burning sodium in 40 gm. of chlorine?

7. Calculate the simplest formula from (a) Hg = 84.92, CI = 15.07; (6) Hg = 73.8, CI = ^6.2.

CHAPTER XII ACroS, SALTS, AND BASES

161. What are acids ? — Hydrochloric acid is an ex- ample of an important class of compounds called acids. The general properties of hydrochloric acid are charac- teristic of the class.

Hydrogen is a constituent of all acids — an essential constituent. The hydrogen of acids can be replaced by certain metals ; and the compound formed by this replace- ment is called a salt. Many compounds have hydrogen as a constituent, but they are not classed as acids unless they form salts by replacement of the hydrogen by a metal. Thus, water and sugar contain hydrogen, but the hydrogen in water does not form a salt by replacement of its hy- drogen by a metal, nor can the hydrogen in sugar be re- placed by a metal. Furthermore, most acids have a sour taste and change the color of a dye called htmus from blue to red. Substances which act thus on blue htmus are said to have an acid reaction.

The presence of acids is often conveniently detected by the litmus test. For example, vinegar, pickles, many fruits, and some wines have a sour taste and turn blue litmus red.

Other common acids are sulphuric (H2SO4), nitric (HNO3), and acetic (HC2H3O2).

162. What are salts? — These compounds are not sour, nor do they, as a rule, have any effect on either blue or red htmus. Substances which have no effect on litmus

142

ACIDS, SALTS, AND BASES 143

are often described as havin"; a neutral reaction. Many- salts have the taste associated with a famiUar member of this class, viz. common salt or sodium chloride ; a few are sour and some are bitter. This class of compounds has many members and their properties are somewhat varied.

Salts invariably have a metal and a non-metal as con- stituents, and most salts also have oxygen. Chlorides are examples of salts which have only a metal and the non- metal chlorine as constituents, e.g. zinc chloride (ZnClo), sodium chloride (XaCl), calcium chloride (CaCb). Ex- amples of salts composed of a metal, a non-metal, and also oxygen are potassium chlorate (KCIO3) and sodium sul- phate (NaoS04) ; these compounds are salts of chloric acid and sulphuric acid respectively.

163. What are bases? — Hydroxides are examples of the class of compounds called bases. Thus, sodium hy- droxide (NaOH) IS a typical base. Solutions of bases turn red Ktmus blue — just the opposite of acids, and are said to have a basic or an alkaline reaction. Solutions of strong bases {e.g. sodium hydroxide and potassium hydroxide) have a slippery feeUng and a biting, caustic taste.

Bases are composed of a metal, oxygen, and hydrogen, e.g. sodium hydroxide (NaOH). The oxygen and hy- drogen are the essential part of a base, just as hydrogen is of an acid. In fact, the oxygen and hydrogen of a base act as a unit in many chemical changes. This group of atoms (OH) is called hydroxyl. And since hydroxy 1 is the root or foundation of a class of compounds, it is called a radical.

164. Neutralization. — Acids, salts, and bases have fundamental chemical relations. When we mix solutions containing weights of an acid and a base proportional to

144 PRACTICAL CHEMISTRY

their molecular weights, the acid and base interact com- pletely. The final solution has none of the characteristic properties of an acid or a base, but it does have the prop- erties of a salt. That is, the acid and base destroy the marked properties of each other, and a salt is produced. The acid and base neutrahze each other. For example, when . hydrochloric acid and sodium hydroxide interact, sodium chloride and water are formed. The equation for the reaction may be written : —

Acid Base Salt Water

HCl + NaOH = NaCl + HoO

Hydrochloric Sodium Sodium Water

Acid Hydroxide Chloride

A chemical change in which an acid and a base neutralize each other and form a salt and water is called neutrali- zation.

Neutralization illustrates double decomposition. In the chemical change just cited both the hydrochloric acid and the sodium hydroxide are decomposed and their parts are recombined in a different way, i.e. sodium chloride and water are the new compounds resulting from the recom- bination.

Neutralization is usually done by means of burettes. These are graduated glass tubes, so marked that any desired portion of the contents can be drawn off by the stop-cock at the lower end (Fig. 70). In using burettes for neutralization, one is filled to the zero (upper) mark with a solution of an acid and the other with a solution of a base — one solution being of known concentration. A measured portion, say 15 cubic centimeters of the base, is drawn off into a beaker, several drops of litmus solution are added, and the acid is slowly dropped in with constant stirring until one drop more shows by the change in color (after thorough stirring) that the right

ACIDS, SALTS, AND BASES

145

proportions of acid and base are present, i.e. that neutralization has occurred. If this solution is evaporated, nothing but a salt will be found in the residue.

If we wish to find the strength of the acid, the volume of acid is read accurately. Knowing the concentration (and volume) of the base solution and the volume of the acid solu- tion, we can calculate the exact weight of the acid needed for the neutralization of the base, and from this weight we can find the strength of the acid solution.

QP=^=^

Fig. 70. — Burettes

165. Another definition of a salt. —

For the present, we may regard salts as compounds formicd from acids and bases by neutralization. That is, the metal of the base unites with the non-metal or non-metallic group of the acid, e.g. Na of NaOH unites with NO3 of HNO3 to form the salt NaNOs. The nature and interrelation of acids, salts, and bases will be further discussed. (See Chapter XVIII.)

166. Naming acids, salts, and bases. — There are three other non-oxygen acids besides hydrochloric acid, viz. hy- drofluoric (HF), hydrobromic (HBr), and hydriodic (HI). The salts corresponding to these four acids end in ide. e.g. chloride, fluoride, bromide, and iodide. Sometimes the compound commonly known as hydrogen sulphide (H2S) is called an acid, and its salts are called sulphides.

Oxygen is a constituent of most acids and salts, and the names of the oxy-acids and oxy-salts are related, especially the suffixes. This relation can be best illustrated by the chlorine acids that contain oxygen. These acids are hypochlorous (HCIO), chlorous (HCIO2), chloric (HCIO3), and perchloric (HCIO4). In forming the names of the corresponding salts the suffix ous becomes ite, while ic be-

146 PRACTICAL CHEMISTRY

comes ate; the prefixes are not changed. Thus, the names of the corresponding sodium salts are sodium hypochlorite, chlorite, chlorate, and perchlorate respectively. In the case of the acids of some elements the body of the name is modified, e.g. sulphuric becomes sulphate (not sulphur- ate!), and phosphoric becomes phosphate (not phosphor- ate!).. Hence we may say briefly, ous becomes ite and ic becomes ate in changing names of acids to the corresponding salts.

Bases are distinguished by placing the name of the metal before the word hydroxide, e.g. sodium hydroxide (NaOH), calcium hydroxide (Ca(0H)2).

167. Conclusion. — More will be learned about acids, bases, and salts, especially in Chapter XVIII and under typical examples of these compounds.

EXERCISES

1. Prepare a summary of this chapter.

2. State some characteristics of (a) acids, (b) bases, (c) salts.

3. How are acids, bases, and salts related?

4. Give the name and formula of three common (c) acids, and

(b) bases, (c) Of five salts.

5. Define and illustrate neutralization.

6. What is hydroxyl?

7. Give the name and formula of the sodium salt of hydrochloric acid. Also of the corresponding salt of potassium, aluminium, lead, silver, antimony, manganese, zinc, and barium.

8. Apply Exercise 7 to (a) nitric acid, (b) nitrous acid, (c) hypo- chlorous acid.

9. Apply Exercise 7 to (a) sulphuric acid, (b) sulphurous acid,

10. Give the name and formula of the hydroxide of the metals enu- merated in Exercise 7.

11. Give the name of these : (a) potassium salt of chloric acid (b) calcium salt of hypophosphorous acid, (c) sodium salt of carbonic acid, (d) lead salt of chromic acid, (c) zinc salt of hydriodic acid, (/) po- tassium salt of perchloric acid, (g) iron salt of hydrochloric acid, (h) so- dium salt of hydrofluoric acid, (i) calcium salt of persulphuric acid,

ACIDS, SALTS, AND BASES 147

{j) potassium salt of hydrobromic acid, (k) calcium salt of hydrofluoric acid, (/) sodium s^lt of hypophosphorous acid.

12. Classify into acids, bases, and salts: KOH, HBr, NH4OH, NaXOs, H3P64, Ag2S04, Ca{OH).,, HI, Pb(0H)2, FeCls, Zn(X03)2, Cu(0H)2, KCIO4, NazCOa, BaS04, BaCl^, A1(0H)3.

PROBLEMS

1. Calculate the per cent of hydrogen in {a) hydrochloric acid, (ft) sulphuric acid, (t) nitric acid.

2. Calculate the per cent of hydroxyl in (a) sodium hydroxide, (h) potassium hydroxide, (c) ammonium hydroxide, (d) calcium hydrox- ide.

3. How many grams of . hydroxyl correspond to (a) 35 gm. of A1(0H)3, (b) 80 gm. of barium hydroxide?

4. What weight of nitric acid is needed to neutralize 27 gm. of the base corresponding to {a) Ca, (b) sodium, (c) K?

5. Two solutions are well mixed. One contained 75 gm. of sul- phuric acid, and the other 75 gm. of sodium hydroxide. What will be the weight of the compounds (other than water) in the final solution?

6. Complete and balance : (a) BaO -\ = Ba(0H)2; (b) XH4I

+ = Agl + (NH4)2S04; (c) Pb(N03)2 + = PbCl2 + .

7. Calculate the formula corresponding to: (a) Ca = 29.41, S = 23.52, O = 4705; (b) Xa = 39.31, CI = 60.68.

8. What weight of the salt is formed in these cases of neutraliza- tion? (a) Hydrochloric acid and 10 gm. of potassium hydroxide; (b) sulphuric acid and 37 gm. of sodium hydroxide.

9. Suppose 37.5 cc. of a hydrochloric acid solution neutralize 30 cc. of a sodium hydroxide solution, and that each cc. of the sodium hydrox- ide solution contains 0.003 S^- of the base. What weight of hydro- chloric acid is contained in the acid solution?

CHAPTER XIII

AMMONIA — AMMONIUM HYDROXIDE — AMMONIUM COMPOUNDS

168. Introduction. — The term ammonia includes both the gas (XH3) and its sokition in water (NH4OH). Some- times the solution is called ammonia water, though its scientific name is ammonium hydroxide. Ammonium hydroxide is an example of a base (163).

169. Formation of ammonia. — When vegetable and animal matter containing nitrogen decays, the nitrogen and hydrogen are usually Hberated as the compound am- monia. The odor of ammonia can be detected near stables. If animal substances containing nitrogen are heated (es- pecially with lime or soda-lime), ammonia is given oE. (The formation of ammonia in this w^ay is a test for com- bined nitrogen. See 114.) Soft coal contains combined nitrogen and hydrogen, and' when the coal is heated, as in making illuminating gas, ammonia is liberated. This is one source of commercial ammonium hydroxide.

170. Preparation of ammonia gas in the laboratory. — Ammonia gas is prepared in the laboratory by heating ammonium chloride w^ith a base, usually moist calcium hydroxide. The equation for the reaction is: —

2NH4CI + Ca(0H)2 = 2NH4OH + CaClo

Ammonium Calcium Ammonium Calcium

Chloride Hydroxide Hydroxide Chloride

The ammonium hydroxide is unstable, especially when

148

AMMONIA — AMMONIUM COMPOUNDS 149

heated, and quickly decomposes into ammonia and water, thus : —

NH4OH = XH3 + H.O

Ammonium Hydroxide Ammonia

Water

The gas is very volatile, and is usually collected by upward displacement, i.e. by allowing the gas to flow upward into a bottle and displace the air (Fig. 71). The solution is prepared by conducting the gas into water.

171. Manufacture of commercial ammonium hydroxide. — The am- monia gas from which commercial ammonium hydroxide is manufactured is obtained mainly from the illuminating gas works. When coal is heated in closed retorts, a mixture of gases is liberated contain- ing ammonia. The am- monia is separated from the other gases by dissolving it out with water (326). This impure solution, which is called ammoniacal liquor or gas liquor, is treated with lime to liberate the ammonia, which is absorbed in tanks containing hydrochloric acid or sulphuric acid. This so- lution upon the addition of a base (e.g. calcium hydroxide) gives up its ammonia, which is dissolved in distilled water, forming thereby the ammonium hydroxide of commerce.

Some ammonia is obtained from the gases Hberated from coke ovens. Ammonia is also manufactured by the direct combination of nitrogen and hvdro^ren fl74. 175).

Fig. 71. — Apparatus for preparing ammonia in the laboratory

I50 PRACTICAL CHEMISTRY

172. Properties of ammonia. — Ammonia gas is color- less. It has an exceedingly pungent odor, and if inhaled suddenly or in large quantities, it brings tears to the eyes and may cause suffocation. It is a light, volatile gas, being about one half (0.59) as heavy as air. A hter of the gas at 0° C. and 760 mm. weighs 0.77 gm. Ammonia gas is easily hquefied — 0° C. and 4.2 atmospheres {i.e. 4 X 760 mm.) being the usual conditions. Liquefied ammonia is often called anhydrous ammonia, because it contains no water.. It boils at —34° C. Hence, if it is exposed to the air or warmed in any way, it changes into the gas, and in so doing absorbs considerable heat. This fact has led to the extensive use of Uquid ammonia in refrigera- tion and in the manufacture of ice (177).

Ammonia gas is very soluble in water, even more so than hydrogen chloride ; and its solubihty can be shown by the " fountain " experiment. (See Fig. 68.) A hter of water at 0° C. dissolves 1148 1. of gas (measured at 0° C. and 760 mm.), while at ordinary temperature i 1. of water dis- solves about 700 1. of gas. This solution of the gas is often called ammonia, though other names, e.g. ammonium hy- droxide and ammonia water are sometimes applied to it (168) ; it gives off the gas freely, when heated, as may be easily discovered by the odor or by the formation of dense white fumes of ammonium chloride (NH4CI) when the solution is exposed to hydrochloric acid (156). The commercial solution called ammonia is lighter than water (its specific gravity being about 0.88) and contains approximately 35 per cent (by weight) of the compound NH3.

173. Chemical conduct of ammonia. — Ammonia gas will not burn in air under ordinary conditions, nor will it support combustion, as the term is usually used ; but if the

AMMONIA — AMMONIUM COMPOUNDS 1 5 1

air is heated or if its proportion of oxygen is increased, a jet of ammonia gas will burn in it with a yellowish flame. When electric sparks are passed through ammonia gas, it decomposes to some extent into nitrogen and hydrogen ; whereas nitrogen and hydrogen form a small amount of ammonia under the same conditions (174. 175).

Ammonia reacts with certain elements. Dried ammonia gas and heated magnesium form magnesium nitride and hydrogen, thus : —

2NH3 + 3Mg = MgsNo + 3H2

Ammonia Magnesium Magnesium Nitride Ammonia and chlorine interact, thus : —

2NH3 + 3CI.2 = No + 6HC1

Ammonia Chlorine Nitrogen Hydrochloric Acid

Ammonia combines directly with water, form^ing ammo- nium hydroxide, thus : —

NH3 + H.2O = NH4OH

Ammonia Water Ammonium Hydroxide

It also combines with certain gases, e.g. hydrogen chloride (HCl) , thereby forming ammonium chloride (NH4CI) . This reaction serves as a test for ammonia gas. (Compare 156.) 174. Synthesis of ammonia from its elements. — Ni- trogen and hydrogen unite if electric sparks are passed through a mixture of these gases. The equation is : — No + 3H2 = 2NH3

The amount of ammonia formed in a given case, however, is only a small per cent of that indicated by the equation. The small yield is due partly to the fact that ammonia itself decomposes into nitrogen and hydrogen. Thus, we may write this equation : —

2NH3 = No + 3H2

152 PRACTICAL CHEMISTRY

If we compare these equations, we see that one is the re- verse of the other. This means that one reaction undoes the work of the other. When the experiment is done in a closed tube, the two reactions proceed at the same time — one reversing the other. Such a complete reaction is called a reversible reaction. The equation for a reversible reaction contains oppositely pointed arrows in place of the equaHty sign. The equation for this reversible re- action is : —

N2 + 3H2 :^ 2NH3

This equation is read : nitrogen and hydrogen react re- versibly to form ammonia.

A reversible reaction under a given set of conditions proceeds to equilibrium. This means that the amounts of the substances involved in both reactions increase or decrease until the quantity of any one substance formed equals the quantity of it which is transformed. In the case of ammonia, the mixture at equilibrium is only about 2 per cent ammonia (and 98 per cent nitrogen and hydro- gen). If the ammonia is removed by adding acid or water to the apparatus, or by some other device, its removal dis- places the equilibrium and the reaction proceeds to com- pletion, i.e. all, or practically all, the nitrogen and hydrogen combine. (See also 196.)

175. Manufacture of ammonia from nitrogen and hy- drogen. — The reaction described in 174 for the manufacture of ammonia proceeds too slowly. To be profitable com- mercially, the reaction must be hastened, that is, its veloc- ity must be increased so that more ammonia will be formed in a given time. Several factors affect the velocity of a reaction, e.g. temperature, pressure, and experimental con- ditions. In manufacturing ammonia, the best conditions

AMMONIA — AMMONIUM COMPOUNDS 153

are a temperature of about 450° C. and a pressure of about 200 atmospheres (i.e. 200 X 760 mm.). This reaction, as well as many others, is further hastened by passing the heated and compressed gases over a catalyst (53) — in this case essentially finely divided iron.

The mixture of ammonia, nitrogen, and hydrogen is passed, still under pressure, through water, which dissolves the ammonia, while the other gases are returned to the apparatus. The ammonia solution is drawn off when it becomes saturated (under ordinary pres- sure and temperature). Ammonia made directly from its elements is called syn- thetic ammonia.

176. Composition of ammonia. — Experi- ments show that ammonia is a compound of nitrogen and hydrogen, especially its synthesis from these elements. By utilizing the fact that ammonia and chlorine react and Hberate nitrogen, the volumetric composition of am- monia can be shown to be nitrogen is to hydro- gen as I to 3. A supplementary experiment Fig. 72. Appara

shows that two volumes of ammonia arc formed tus for deter by the union of one volume of nitrogen and three volumes of hydrogen.

The volumetric composition of ammonia may be expressed thus : —

N2 -f- 3H2 = 2NH3

I vol. of 3 vols, of 2 vols, of

Nitrogen Hydrogen Ammonia

mining the volu- metric composi- tion of ammonia gas by the inter- action of ammo- nium hydro.xide and chlorine

The simple volumetric relation of these three gases is a good example of Gay-Lussac's law (93, 157).

In demonstrating the volumetric composition of ammonia, a tube of known volume (Fig. 72) filled with chlorine is provided with a funnel through which concentrated ammonium hydroxide is slowly dropped into the chlorine, until the reaction ceases. After the ex-

154

PRACTICAL CHEMISTRY

cess of ammonium hydroxide is neutralized with sulphuric acid, the volume of nitrogen left is found to be one third of the original volume of chlorine. Now hydrogen and chlorine combine in equal volumes (157). Hence the volume of hydrogen withdrawn from the ammonia must be equal to the original volume of chlorine. But this volume is three times the volume of the nitrogen, therefore there must be three times as much hydrogen as nitrogen in ammonia gas.

177. Ammonia as a refrigerant. — The use of ammonia in producing low temperatures depends upon the fact that liquefied ammonia (not ordinary ammonia solution) changes rapidly into a gas when its pressure is reduced, and in so do- ing absorbs heat from the surrounding air or liquid. Hence,

Fig. 73. — Apparatus for using liquefied ammonia to produce low temperatures

if liquefied ammonia is allowed to flow through a pipe im- mersed in a solution of sodium chloride or calcium chloride (technically called a brine), the ammonia evaporates in the pipe and cools the brine, which may be used directly as a refrigerant or for making ice. In some cold storage plants, packing houses, and sugar refineries, this cold brine is circulated through pipes placed in the storage rooms where a low temperature is desired.

The construction and general operation of an ice-making plant is shown in Fig. 73. Liquefied ammonia is forced from a tank into a series of pipes which are submerged in a large vat A nearly filled with

AMMONIA — AMMONIUM COMPOUNDS 155

brine. Metal cans containing pure water to be frozen are immersed in the brine, which is kept below the freezing point of water by rapid evaporation of the ammonia in the pipes. After several hours the water in the cans is frozen into cakes of ice. As fast as the ammonia gas forms in the pipes, it is removed by exhaust pumps (£) into another set of pipes C, where it is condensed into liquefied ammonia and con- ducted through D into the other set of pipes ready for renewed use. In cold storage plants the cold brine is circulated through pipes in the various rooms {B).

178. Ammonium hydroxide. — When ammonia gas is passed into water, the ammonia combines with the water to some extent and forms a solution of an unstable com- pound having the composition represented by the formula NH4OH. This compound is ammonium hydroxide. Am- monium hydroxide is a base (163). Like other members of this class of substances it turns litmus blue. Concen- trated solutions have a slippery feehng. It also neutral- izes acids, thus : —

NH4OH -1- HCl = NH4CI + H2O

Ammonium Hydrochloric Ammonium Water

Hydroxide Acid Chloride

Ammonium hydroxide is widely used as a cleansing agent (especially for the removal of grease), and as a restorative in cases of fainting or of inhaling irritating gases; large quantities are consumed in dyeing and calico printing, and in the manufacture of dyestuffs, sodium carbonate and bicarbonate, and ammonium compounds.

179. Ammonium compounds contain a group of atoms which acts chemically like an atom of a metal, especially the metals sodium and potassium. This group is called ammonium, and its formula is NH4. Ammonium has never been isolated from its compounds. Ammonium, like hydroxyl (163), is called a radical, because it is the

156 PRACTICAL CHEMISTRY

root or foundation of a series of compounds and in chemical changes passes from one ammonium compound to another. Ammonium compounds decompose when heated with an alkaU, such as sodium hydroxide or moist Hme, i.e. calcium hydroxide (170) ; ammonia gas is the conspicuous product. This reaction is a test for ammonium com- pounds.

180. Ammonium chloride. — There are many ammo- nium salts. Thus, ammonium chloride (NH4CI) is the salt formed by the neutralization of ammonium hydroxide (a base) by hydrochloric acid (see 164). It is manufactured by passing ammonia (obtained from the ammoniacal liquor in gas works) into hydrochloric acid. The crude product is often called " muriate of ammonia " to indicate its re- lation to muriatic acid (the commercial name of hydro- chloric acid). It is used in Leclanche batteries, as an in- gredient of soldering fluids and of fertilizers.

The crude salt is purified by heating it gently in a large iron pot with a dome-shaped cover ; the ammonium chloride volatilizes easUy and then crystallizes in the pure state as a fibrous mass on the in- side of the cover, whereas the impurities remain behind in the vessel. This process of purification is called sublimation. The product is a sublimate. Sublimed ammonium chloride is known as sal ammoniac.

181. Other ammonium salts. — Ammonium sulphate ((NH4)2S04) is made by passing ammonia gas into sulphuric acid, or by adding ammonium hydroxide to the acid, thus : —

2NH4OH + H0SO4 = (NH4)2S04 + HoO

Ammonium Hydroxide Ammonium Sulphate

The commercial salt is a grayish or yellowish solid and is obtained in large quantities as a by-product of coal gas manufacture. It is used as an ingredient of fertilizers, since it is a cheap, soluble salt containing considerable nitrogen (116). Ammonium nitrate (NH4NO3) is made by passing ammonia into nitric acid, or by allowing ammonia gas and the vapor of nitric acid to mingle, thus : —

AMMONIA — AMMONIUM COMPOUNDS 157 NH3 + HXOx = NH4NO3

Ammonia Nitric Acid Ammonium Nitrate

When gently heated it decomposes into nitrous oxide (N2O) and water, and is used in the preparation of nitrous oxide. It is also used as an ingredient of certain explosives because it decomposes readily into volatile products. Ammonium carbonate ((NH4)2C03) is used in some kinds of baking powder, to scour wool, as a medicine, and in certain smelling salts (since it gives off ammonia readily).

EXERCISES

1. Prepare a summary of ammonia.

2. How is ammonium hydroxide manufactured from (a) gas liquor, (b) nitrogen and hydrogen?

3. State the conspicuous properties of ammonia gas.

4. Describe the synthesis of ammonia.

5. Define and illustrate (by ammonia) (a) reversible reaction

(b) equilibrium, (c) displacement of equilibrium, (d) catalyst.

6. State these reactions in the form of equations: (a) Prepar- ation of ammonium hydroxide from ammonium chloride and calcium hy- droxide, (b) Decomposition of ammonium hydroxide into ammonia and water, (c) Nitrogen and hydrogen react reversibly to form am- monia, (d) One volume of nitrogen and three volumes of hydrogen form two volumes of ammonia.

7. What is (a) liquid ammonia, (b) anhydrous amrronia, (c) lique- fied amrfionia, {d) synthetic ammonia, (c) ammonia water?

8. Describe the manufacture of ice by liquid ammonia.

9. What is the volumetric composition of ammonia gas? How is it found?

10. State the test for (a) ammonia, (b) ammonium compounds,

11. Topics for home study: (a) A cold storage plant, {b) Gay- Lussac's law. (c) Uses of ammonia, {d) Uses of ammonium com- pounds, (c) Equilibrium.

PROBLEMS

1. How many grams of ammonia (XH3) can be obtained from i kg. of ammonium chloride and sufficient calcium hydroxide?

2. A pupil prepared five 250 cc. bottles of ammonia gas at 21° C. and 755 mm. What weight of materials interacted?

3. What weight of ammonium chloride (95 per cent pure) is needed for the preparation of 60 gm. of NH3? Of 60 1. at 22° C. and 767 mm.?

158 PRACTICAL CHEMISTRY

4. What is the weight of 32 1. of ammonia gas at 20° C. and 763 mm.? What volume will 32 gm. of ammonia gas occupy at the same temperature and pressure?

5. To what weight and what volume of NH3 are 25 gm. of am- monium chloride equivalent (at standard conditions) ?

6. What volume of ammonia gas will be liberated by the action of any base on 75 gm. of ammonium sulphate (at standard conditions)?

7. How many liters of ammonia are formed by the complete inter- action of 9 1. of hydrogen and sufficient nitrogen?

8. How many grams of ammonium chloride can be made from am- monium hydroxide and 100 gm. of the necessary acid?

9. What formula corresponds to N = 26.17, H = 7.48, CI = 66.35?

10. Complete and balance : (a) NH4OH H = (XH4)2S04

+ ; (b) NH3 + = N2 + HCl.

CHAPTER XIV

NITRIC ACm — NITRATES — NITROGEN OXIDES

C=^3

182. How nitric acid is formed. — Nitric acid (HNO3) is formed in small quantities when electric sparks are passed through air. Hence nitric acid or its salts can be detected in the atmosphere after a thunderstorm. This chemical change is ap- plied on a commercial scale in Norway (196). In another chemical pro- cess ammonia is oxidized (197).

183. Preparation of nitric acid. — Nitric acid is prepared in the labora- tory by heating concen- Fig trated sulphuric acid with a nitrate, usually sodium nitrate (NaNOs).

About equal weights of sodium nitrate and concentrated sulphuric acid are put into the glass retort (Fig. 74), and gently heated. The nitric acid is much more volatile than the sulphuric acid and distills into the receiver {e.g. a large test tube), which is kept cool by water, ice, or moist paper.

The chemical change at a low temperature is expressed by the equation : —

NaN03 + H.2SO4 = HNO3 + HNaS04

Sodium Sulphuric Nitric .\cid Sodium

Nitrate Acid Acid Sulphate

159

74. — .Apparatus for preparing nitric acid in the laboratory

i6o

PRACTICAL CHEMISTRY

But if the temperature is high and an excess of sodium ni- trate is used, the equation is : —

2NaN03 + H2SO4 = 2HNO3 + Na2S04

At a high temperature part of the nitric acid decomposes ; hence excessive heat is usually avoided.

184. Manufacture of nitric acid. — Nitric acid is man- ufactured by a process like that used in the laboratory. A sketch of the apparatus is shown in Fig. 75.

Fig.

Sketch of the apparatus for manufacturing nitric acid

acid and sodium nitrate are heated in the cast iron retort, which is connected with glass or stoneware tubes in which the vapor is condensed by a current of cold water ; the tubes are arranged so that the nitric acid first condensed runs into a reservoir, while the vapors pass up a tower, where they are dissolved by descending water and flow- out at the bottom as dilute acid. (See 196, 197.)

185. Properties of nitric acid. — Pure nitric acid is a colorless liquid, but the commercial acid is yellow or red- dish, due to dissolved nitrogen oxides. The acid that has been exposed to the sunlight is often yellow or brown, and if the light is intense, a brownish gas may often be seen in bottles of the acid. It is somewhat volatile, and the vapor dissolves readily in water ; hence the acid forms irritating

NITRIC ACID — NITRATES t6i

fumes when exposed to air, especially moist air. (See 155.)

Nitric acid mixes with water in all proportions. Com- mercial concentrated nitric acid contains about 68 per cent of the compound HXO.j, the rest being water. Such an acid has a specific gravity of 1.41 and boils at about 120° C.

186. Chemical conduct of nitric acid. — Nitric acid is sour, turns blue litmus red, and forms salts — the nitrates. It is an unstable substance, and decomposes readily; among the decomposition products is a brown gas, nitrogen dioxide (NO2), which causes the yellow or brown color referred to above (185). If strongly heated, the decom- position is rapid ; the equation is : —

4HNO3 = 4XO0 + O2 + HoO

Nitric Acid Nitrogen Dioxide Oxygen Water

Nitric acid is a very corrosive substance and reacts readily with many substances. With organic substances Hke hair, feathers, w^ooL silk, finger nails, and skin it forms a yellow compound. Hence nitric acid stains the skin, and often the clothing, yellow. The concentrated acid causes serious burns and should not be spilled on the hands or face. With other organic conipounds it forms explosives, such as nitroglycerin and nitrocellulose.

One of the decomposition products of nitric acid is oxy- gen. Hence nitric acid is an oxidizing agent. Usually the oxygen is not liberated as a gas, but oxidizes whatever oxidizable substance is present. In order to emphasize the fact that oxidizing oxygen is available, the equation is written thus : —

2HNO3 = 3O + 2NO + HoO

Nitric Acid Nascent Oxygen Nitric Oxide Water

i62 PRACTICAL CHEMISTRY

Thus, charcoal burns briUiantly in the hot acid, while straw, sawdust, hair, and similar substances are charred and even inflamed by it; some organic compounds, w^hen heated with nitric acid, are completely decomposed into carbon dioxide and water. In the mixture of concentrated nitric and hydrochloric acids called aqua regia, nitric acid acts as an oxidizing agent (159). The nitric acid oxidizes the hydrochloric acid to water and leaves the chlorine in the nascent condition. (Compare end of 149.) The equation is usually written : —

HN03

+ 3HCI =

2H2O + 2CI + XOCl

Nitric

Hydrochloric

Water Nascent Nitrosyl

Acid

Chlorine Chloride

Finally, it interacts readily and often violently with metals, metalHc oxides, and hydroxides. (Compare 158.) The products of these reactions vary, the essential ones being nitrates and nitrogen oxides (189).

187. Uses of nitric acid. — Nitric acid is one of the com- mon laboratory acids. Large quantities are used in the manufacture of nitrates, dyestuffs, sulphuric acid, and explosives, and in etching copper plates.

188. Nitrates. — Nitric acid forms salts called nitrates. They are prepared by the methods usually used for salts, i.e. the interaction of nitric acid and metals or metalhc oxides and the neutralization of hydroxides by nitric acid.

Many nitrates "are white sohds ; those of copper, nickel, and cobalt are blue, green, and dark red respectively. The nitrates of most metals are soluble in water. Their solu- tions are frequently used in the laboratory. The solids behave in various ways when heated. Equations illus- trating typical reactions are : —

NITRIC ACID — NITRATES 163

2NaN03 = 2NaN02 + O2 ;

Sodium Nitrate Sodium Nitrite

2CU(X03)-: = 2CuO + 4NO2 + O2 Copper Nitrate Copper Oxide Nitrogen Oxide

Since many nitrates, when heated, give up oxygen, they are powerful oxidizing agents. Thus, when potassium nitrate is dropped on hot charcoal, the charcoal burns vigor- ously. This kind of chemical action is called deflagra- tion.

Nitrates are also formed in the soil by the action of bac- teria on complex nitrogen compounds. This process, w^hich is called nitrification, is slow ; formerly this was the sole source of the potassium nitrate needed for gunpowder.

189. The interaction of nitric acid and metals. — This action is exceedingly vigorous; for this reason, probably, the alchemists called the acid aqua fortis — strong water. The products of the reaction vary with the metal, the con- centration of the acid, and the temperature. Hydrogen as a rule is not liberated so that it can be collected, for it is oxidized at once to water by the nitric acid, whereas the nitric acid is reduced to nitrogen compounds.

The interaction of nitric acid and copper will serve as an example of the common reactions. When moderately dilute nitric acid is poured on copper, a reddish brown gas is given off, and the nquid turns blue, owing to dissolved copper nitrate. The equation for the reaction is: —

3CU -f 8HNO3 = 3Cu(N03)2 + 2NO + 4H2O

Copper Nitric Copper Nitric Water

Acid Nitrate Oxide

This equation is made up of three equations, and the

i64 PRACTICAL CHEMISTRY

complete form conceals the way in which the reactions take place. The nitric acid first decomposes, thus : —

(i) 2HNO3 = 30 + 2NO + HoO

Nitric Acid Nascent Oxygen Nitric Oxide Water

The nascent oxygen (149) next oxidizes the copper, thus : —

(2) 3Cu + 30 = sCuO

Copper Nascent Oxygen Copper Oxide

The copper oxide then reacts with the nitric acid, thus : —

(3) 3CUO + 6HNO3 = 3Cu(N03)2 + 3H2O

Copper Oxide Nitric Acid Copper Nitrate Water

Since 3O is formed in (i) and used in (2) and 3CUO like- wise in (2) and (3), these two terms should not appear in the complete equation ; the other terms make up the com- plete equation.

Nitric oxide is represented as a product of the interaction of nitric acid and copper. If the reaction takes place in an open vessel, the nitric oxide, which is a colorless gas, com- bines with oxygen and forms the reddish brown nitrogen dioxide gas. The equation is : —

2NO + Oo = 2NO2

Nitric Oxide Oxygen Nitrogen Dioxide

Hence we often speak of nitrogen dioxide as a product of the interaction of nitric acid and metals, though it is usu- ally a secondary product.

190. The test for a nitrate (and of course for nitric acid) is not, as customary, made by producing a precipitate. It is a color test, and is made as follows : Add to the so- lution of the nitrate or the nitric acid in a test tube an equal volume of ferrous sulphate solution (freshly prepared from clean ferrous sulphate and cold water) ; mix well. In-

NITRIC ACID — NITRATES 165

dine the test tube, and pour concentrated sulphuric acid cautiously down the side of the test tube. A dark brown layer appears where the two liquids meet (Fig. 76), owing to the formation of a brown unstable com- pound which has the composition (approx- imately) 3FeS04.2NO.

191. Nitrous acid (HNO2) is not easily obtained in the free state, owing to its instability, but its salts — the nitrites — are well known. Potassium _ nitrite (KNO2) and sodium nitrite (NaNOo) are ^. ^ ^ formed by removing part of the oxygen from the T ' '. . . , corresponding nitrate by heating alone (188) or ^^^ nitrates with lead. Nitrites give off yellow-brown fumes

(NOo) readily when treated with sulphuric acid, and are thus easily distinguished from nitrates.

192. Nitrogen oxides. — There are live nitrogen oxides. The three important ones are nitrous oxide (N2O), nitric oxide (NO), and nitrogen dioxide (NO2) ; there is also an oxide called nitrogen tetroxide (N2O4), which is a special form of nitrogen dioxide.

193. Nitrous oxide (N2O) is prepared by heating am- monium nitrate. The equation for the reaction is : —

NH4NO3 = N2O + 2H2O

Ammonium Nitrate Nitrous O.xide Water

This colorless gas has a faint but pleasant odor. It is sol- uble in water, and the solution has a sweet taste. It is easily liquefied by reducing the temperature and applying pressure, and is often used in this form to furnish the gas. The gas does not burn, but it supports the combustion of many well-burning substances, though not so vigor- ously as oxygen does. Thus, sulphur, unless well ignited, will not burn in nitrous oxide. In its power to support

1 66 PRACTICAL CHEMISTRY

combustion it resembles oxygen. It is distinguished from oxygen by its failure to form brown fumes (NO2) when mixed with nitric oxide (194).

The most striking property of nitrous oxide is its efTect on the human system. If inhaled for a short time, it causes more or less nervous excitement, often manifested by laughter, and on this account the gas was called " laughing gas" by Davy, who first studied its properties in 1799. If breathed in large quantities, it produces temporary un- consciousness and insensibiUty to pain. The gas, mixed with a small proportion of air or oxygen, is often used as an anaesthetic in dentistry.

194. Nitric oxide (NO) is usually prepared by the inter- action of copper and dilute nitric acid (sp. gr. 1.2). The complete equation (189) for the reaction is : —

3CU + 8HNO3 = 2NO + 3Cu(N03)2 + 4H2O

Copper Nitric Nitric Copper Water

Acid Oxide Nitrate

Nitric oxide is a colorless gas. It is a little heavier than air and only sHghtly soluble in water. Upon exposure to air, it combines at once with the oxygen, forming reddish brown fumes of nitrogen dioxide — a striking change. The equation for this reaction is : —

2NO + 0-2 = 2NO2

Nitric Oxide Oxygen Nitrogen Dioxide

This property distinguishes nitric oxide from all other gases. It does not burn, nor support combustion, unless the burning substance (e.g. phosphorus or sodium) introduced is hot enough to decompose the gas into nitrogen and oxygen, and then, the liberated oxygen assists the combustion.

195. Nitrogen dioxide (NO2) is the reddish brown gas formed by the direct combination of nitric oxide and oxy-

NITRIC ACID — NITRATES 167

gen (189 end). It is also produced by heating certain

nitrates.

Thus : —

2Pb(N03)2 = 4NO2 + 2PbO + O2 Lead Nitrate Nitrogen Dioxide Lead Oxide Oxygen

The fumes of nitrogen cUoxide usually appear when ni- tric acid and metals interact, but, as stated in 194, the nitrogen dioxide is produced by a second reaction, viz. the combination of nitric oxide with the oxygen of the air. Nitrogen dioxide has a disagreeable odor, and it is poi- sonous if breathed in moderate quantities. It interacts with water and yields under ordinary conditions nitric oxide and nitric acid, thus : —

3NO2 + H2O = 2HNO3 + NO

Nitrogen Dioxide Water Nitric Acid Nitric Oxide

(See 196.) It also dissolves in concentrated nitric acid, forming fuming nitric acid, which is an oxidizing agent.

When the reddish brown gas is cooled, it gradually loses color and at about 26° C. becomes a yellow gas, which has the composition represented by the formula N2O4 and is called nitrogen tetroxide. Upon heating nitrogen tetroxide, the brown gas reappears, and at about 140° C. the gas is wholly nitrogen dioxide. Above 140° C. the brown color fades, owing to the decomposition of nitrogen dioxide into nitric oxide and oxygen. At ordinary temperatures the brown gas is a mixture of the two oxides.

A simple demonstration of the relation between nitrogen diox- ide and tetroxide is readily made by collecting some nitrogen dioxide in a glass tube, closing the tube, and immersing the lower half in ice water. The gas in the lower part becomes yellow-brown (NO4) whereas in the upper part it remains reddish brown (NO2).

196. Manufacture of nitric acid from nitrogen oxides. — If electric sparks are passed through air, nitrogen and oxy- gen unite, thus : —

i68

PRACTICAL CHEMISTRY

(i) No + O2 = 2NO

But the nitric oxide combines at once with oxygen, and forms nitrogen dioxide, thus : —

(2) 2NO + O2 = 2NO2

When nitrogen dioxide is added to water, nitric acid and nitric oxide are produced, thus : —

(3) 3NO2 + H2O = 2HNO3 + NO

These reactions are the basis of modern processes of making nitric acid from nitrogen oxides and water.

Air is used in one process as a source of the nitrogen oxides. Two unusual conditions must be fulfilled in this

L_A

Fig. 77. — Sketch of the apparatus for the manufacture of nitric acid from nitrogen, oxygen, and water

process. First, nitrogen and oxygen must be heated to a very high temperature (about 3000° C.) before they will unite to an appreciable extent. Second, the mixture of gases resulting from reaction (i) must be cooled very quickly. Reaction (i) is reversible, thus: —

N2 + O2 :;=±: 2NO The maximum quantity (only about 5 per cent) of nitric oxide is obtained at about 3000° C. At this temperature equilibrium is reached, i.e. the relative proportions remain unchanged. (Compare 174.) But just as soon as the tern-

NITRIC ACID — NITRATES

169

peraturc becomes lower, the reverse reaction (right to left) begins and the quantity of nitric oxide rapidly decreases. However, if the mixture is cooled very quickly, enough nitric oxide is left for the reaction with water.

J 3R

/til

-Hi'

Fig. 7

spread out (left) and end view (right)

A sketch of the apparatus is shown in Fig. 77 and certain parts are shown in detail in Fig. 78. Air is blown (by A) into the electric fur- nace B. Here it is raised to the proper temperature by passing through an electric arc. In order to provide a large heating surface, the arc is spread out by a magnet into a disk six feet in diameter (Fig. 78, left). The electrodes usually are hollow and are kept cool by running water; graphite electrodes are also used. An end view (Fig. 78, ri^ht) shows the magnets and edge of the disk. The hot gases containing the nitric oxide from the furnace are suddenly cooled in C, pass through boiler D into the oxidizing chamber E, where nitro- gen dioxide is formed, and then into the tower F filled with tiles over which water trickles, where the nitric acid is produced. The dilute nitric acid is concentrated or converted (by limestone or lime) into calcium nitrate. The latter is used as a fertilizer, either alone or mixed with lime.

197. Manufacture of nitric acid from ammonia. — In another process, ammonia is the starting point. A heated mixture of ammonia and air is passed into lire-brick- lined chambers containing a catalyst — usually platinum in the form of gauze. At about 600° C, the reaction, ex- pressed by the following equation, takes place : — 4NH3 + 5O2 = 4NO + 6H2O

I70 PRACTICAL CHEMISTRY

The nitric oxide is cooled, mixed with air, and then, as in the other process, passed into towers through which water trickles, where the nitric acid is produced, thus : —

4NO + 2H2O + 3O2 = 4HNO3

EXERCISES

1. How is nitric acid prepared in the laboratory? State the two equations.

2. Describe the older process of manufacturing nitric acid.

3. Summarize the properties of nitric acid.

4. Explain: (a) nitric acid is an oxidizing agent ; (b) nitric acid is unstable.

5. What is the test for (a) nitric acid, (b) a nitrate, (c) nitric oxide, (d) nitrous oxide, (e) nitrogen dioxide, (/) a nitrite?

6. Describe the interaction of nitric acid and copper and state the equations.

7. Summarize briefly the properties of nitrous oxide.

8. Compare nitric oxide and nitrogen dioxide.

9. How is nitrogen dioxide prepared?

10. State the relation of nitrogen dioxide and nitrogen tetroxide.

11. Describe fully the manufacture of nitric acid from (a) nitrogen oxides and (b) ammonia.

12. Interpret the equation N2 + O2 ^^2_ 2NO.

13. Complete and balance : {a) CuCOs H = Cu(N03)2 H

+ ; (b) HNO3 + C = NO2 + + CO2.

PROBLEMS

1. Calculate the percentage composition of (a) nitric acid, {b) po- tassium nitrite, (c) sodium nitrate.

2. Show how nitrous oxide, nitric oxide, and nitrogen dioxide illus- trate the law of multiple proportions.

3. Calculate the formula and give the name of the compound cor- responding to (a) O = 76.19, H = 1.58, N = 22.22; (b) O = 47-52, K = 38.61, N = 13.86.

4. If the specific gravity of a sample of nitric acid is 1.522, (a) what will 100 cc. weigh; (b) what volume must be taken to weigh 100 gm. ?

5. How many liters of oxygen are needed to change 6 1. of nitric oxide to nitrogen dioxide?

6. How much sodium nitrate is needed to form a ion of commercial nitric acid? (See end of 185.)

CHAPTER XV

MOLECULAR WEIGHTS AND FORMULAS

198. Introduction. — Formulas and molecular weights have been used freely in the foregoing pages. In this chapter we shall consider the methods by which molecular weights and formulas are found by experiment.

199. Gay-Lussac's law of gas volumes. — We have seen several times that the volumes of gases involved in a chemical change can be expressed by small whole num- bers (92, 93, 157, 176). These results may be summarized in a: —

Table of the Combixatiox of Gases by Voli':^ie

Volumes of Combining Gases

2 volumes of hydrogerx I volume of oxygen

I volume of chlorine I volume of hydrogen

3 volumes of hydrogen

1 volume of nitrogen

2 volumes of nitrogen I volume of oxygen

I volume of nitrogen I volume of oxygen

Volumes of Gaseous Product

volumes of water vapor

•olumes of hydrogen chloride

2 volumes of ammonia

2 volumes of nitrous oxide

2 volumes of nitric oxide

It is clear from the above table that the volumes of the gases can be expressed by small whole numbers. This

171

172 PRACTICAL CHEMISTRY

simple relation is true of all gas reactions, and, as we found in studying the volumetric composition of water (92, 93), it may be stated as Gay-Lussac's law, thus : —

In a chemical change the volumes of the gases can he expressed by small whole numbers.

This law applies only to gases. For example, carbon is not a gas, though it is frequently involved in reactions with gases, e.g. : —

Carbon + Oxygen = Carbon Dioxide I vol. I vol.

So we omit carbon as far as volume is concerned, and say carbon unites with I volume of oxygen (O2) to. form i volume of carbon dioxide (CO.2).

200. Avogadro's theory. — In 181 1 the Italian physicist Avogadro proposed an explanation of the simple numerical relation of gas volumes. It is usually called Avogadro's theory and may be stated thus : —

Equal volumes of gases under like conditions of temperature and pressure contain the same number of molecules.

This theory means that a liter of oxygen contains just as many molecules as a liter of hydrogen, nitric oxide, or any other gas, if the temperature and pressure conditions are not altered. The actual number of molecules in the volume of the gas is not important, but the assumption that the number — whatever it may be — is the same in equal volumes is very important in chemistry.

201. How Avogadro s theory is used in finding relative molecular weights. — By means of Avogadro's theory we can find the relative molecular weights of gases. Let us consider carbon dioxide and oxygen. A liter of carbon dioxide weighs 1.98 gm. and a liter of oxygen 1.43 gm. at 0° C. and 760 mm. Therefore the weight of a liter of car- bon dioxide is (approximately) 1.38 times that of -^ liter of oxygen. Since a liter of each gas contains the same number

MOLECULAR WEIGHTS AND FORMULAS 173

of molecules, the weight of the carbon dioxide molecules is (approximately) 1.38 times the weight of the oxygen mole- cules. It is evident, then, that if we weigh equal volumes of gases (under like conditions of temperature and pres- sure), we obtain weights which are in the same relation as the weights of single molecules.

202. How we find the approximate molecular weight of a gas. — In 201 we found that carbon dioxide molecules weigh (approximately) 1.38 times the weight of oxygen molecules. Obviously, we could calculate the (approximate) molecular weight of carbon dioxide if we knew the molec- ular weight of oxy^gen. The molecular weight of oxygen is 32. (See next paragraph.) We now calculate the (approximate) molecular weight of carbon dioxide by multiplying 32 by 1.38, i.e. 32 X 1.38= 44.16. Therefore 44.16 is the (approximate) molecular weight of carbon dioxide.

The molecular weight of oxygen is 32 for two reasons. First, a molecule of oxygen contains 2 atoms, as will be shown in 203. Second, an atom of oxygen weighs 16, be- cause this number has been adopted by chemists as the standard atomic weight (215). Hence the molecular weight of oxygen is 32 {i.e. 2 X 16).

This method of finding molecular weights is called the vapor density method. The steps are (i) find the vapor density referred to oxygen, and (2) multiply this value by 32. The expression vapor density referred to oxygen means the number found by dividing the weight of a given volume of a gas or vapor by the weight of an equal volume of oxygen (measured at the same temperature and pressure). Thus, in the example given above the num])er 1.38 is the vapor density of carbon dioxide {i.e. i.q8 ^ 1.43^*

One method of determining the approximate molecular

T 74 PRACTICAL CHEMISTRY

weight of a gas or a readily volatilized substance is now clear, viz. find the vapor density on the oxygen basis and multiply this value by 32, or

Molecular Weight = Vapor Density referred to Oxygen X 32

Molecular weights determined by this method are ap- proximate. That is, they do not exactly equal the sum of th€ weights of the atom^s in a molecule, i.e. the so-called theoretical molecular weight. The difference is due mainly to slight but unavoidable errors in the process. For ex- ample, the approximate molecular weight of carbon dioxide is 44.16, whereas the exact molecular weight is 44.00 (i.e. 12 + (2 X 16)). The slight difference in this, and other cases, does not affect the validity and use of the weights found by this method.

Some substances cannot be vaporized without decomposition. The molecular weights of such substances cannot, of course, be found by the vapor density method. If a substance dissolves without decomposition, its molecular weight can be determined by an appro- priate method (244).

No experimental method, however, has been devised for deter- mining the molecular weight of a substance in the soHd state {i.e. not dissolved or vaporized) ; it is customary to assume that the molec- ular weight of such substances is the sum of the atomic weights in the simplest formula (109 and last paragraph in 206).

203. How Gay-Lussac's law and Avogadro's theory are used to find the number of atoms in a molecule. — A molecule of oxygen contains two atoms. This conclusion is based mainly on the following argument : When oxy- gen and nitrogen combine to form nitric oxide, the volumes used and produced can be expressed thus : —

Oxygen -f Nitrogen = Nitric Oxide

I vol. I vol. 2 vols.

Now according to Avogadro's theory, equal volumes of

MOLECULAR WEIGHTS AND FORMULAS 175

oxygen and of nitrogen contain the same num])er of mole- cules, while the two volumes of nitric oxide contain twice this number of molecules. We do not know the actual number of molecules in any of the volumes. Suppose, however, there are 1000 molecules of oxygen ; then by Avogadro's theory, there are 1000 molecules of nitrogen and 2000 molecules of nitric oxide, thus : —

Oxygen + Nitrogen = Nitric Oxide

1000 molecules 1000 molecules 2000 molecules

Every molecule of nitric oxide must contain at least one atom of oxygen ; and the 2000 molecules must contain at least 2000 atoms of oxygen. But these 2000 atoms of oxygen were provided by the 1000 molecules of oxygen. Therefore, each molecule of oxygen must contain at least two atoms of oxygen.

By a similar argument it can be shown that a molecule of nitrogen contains at least two atoms. So also it can be shown in the same way that a molecule of the common elementary gases contains at least two atoms.

Furthermore, experiment shows that there is no reaction in which a given volume of the elementary gases, ox}^gen, nitrogen, hydrogen, and chlorine, provides material for more than two volumes of the gaseous product. This means there is no reaction in which a molecule of these gases is divided into more than two parts. And so we conclude that the molecule of these gases contains only two atoms. Hence we write their formulas Oo, N-j, Ho, and CI2. It is clear now why we used these formulas in preceding sections for molecules of the gases.

204. What is a mole? — A mole of a substance is the number of grams numerically equal to its molecular weight. It is sometimes defined as a gram-molecular weight. Thus,

176

PRACTICAL CHEMISTRY

a mole of carbon dioxide is 44 grams because the molec- ular weight is the mmiber 44. Similarly, a mole of oxy- gen is 32 grams, of carbon monoxide is 28 grams, and of nitric oxide is 30 grams.

The volume occupied by one mole of oxygen (at 0° C.

and 760 mm.) is 22.4 liters.

Suppose we construct a box holding 22.4 liters (Fig. 79) and fill it with oxygen (at 0° C and 760 .mm.), the gas will weigh 32 grams. This must be so. One liter of oxygen weighs 1.43 grams. Therefore the volume oc- cupied by 32 grams of oxygen will be 32 ^ 1.43,

22.4 -J LITERS

Fig. 79. — A box holding i mole or 22.4 liters of a gas (actual length of one edge is 28.2 cm.)

or 22.4 liters. Simple division shows that i mole of a gas occupies 22.4 hters (in round numbers), e.g. carbon mon- oxide, 28 -^ 1.25 = 22.4.

This volume (22.4 liters) is sometimes called the gram- molecular volume, since it is the volume of the gram-molec- ular weight. It is also called a molar volume, because it is the volume of one mole. So it is clear that i mole of a gas and 22.4 Hters of the same gas are equal — one being the weight and the other the volume of the same mass of gas.

205. Calculation of molecular weight from a mole. — Suppose we construct several cubical boxes each holding 22.4 liters (Fig. 80), weigh each box, fill each with a gas (at 0° C. and 760 mm.), and weigh again. The increase in each case is the weight of one mole of the several gases, i.e. it is numerically equal to the molecular weight. If the gases used, for example, were nitric oxide, carbon mon- oxide, hydrogen chloride, and ammonia, the numbers ob-

MOLECULAR WEIGHTS AND FORMULAS 177

taincd would be 30, 28, 36.5, and 17. And these numbers are the respective molecular weights of these gases.

Hence to find the molecular weight of a gas by the mole method, we fmd the weight of 22.4 Uters of the gas. It is not necessary to use such a large volume in the actual

22.4 Liters

22.4

Liters

22.4 Liters

Fig. 80.

n c I)

A mole of difTerent gases occupies 22.4 liters. /I = 30 gm. NO, 5 = 28 gm. CO, C = 36.5 gm. HCl, D = 17 gm. XH3

experiment. We simply find the weight of any convenient volume and then calculate the weight of 22.4 hters. An example will make this clear. If 1.5 liters of carbon mon- oxide (at 0° C. and 760 mm.) weigh 1.88 grams, what is the molecular weight of the gas? The weight of 22.4 liters is found thus: 1.88 -^ 1.5 = 1.25; 1.25 X 22.4 = 28. Therefore 28 is the molecular weight of carbon monoxide.

206. Correct formulas of compounds. — A formula of a compound expresses composition. That is, it repre- sents by means of symbols the kind and number of atom.s in a molecule.

In 109 it was shown that the simplest formula of a com- pound can be calculated from the percentage composition, viz. by dividing the per cent of each element in the com- pound by the atomic weight, and then, if necessary, re- ducing the quotients to the smallest whole numbers. A formula thus calculated expresses in the simplest chem- ical way the proportions of the different elements in a com- pound. But it may not be its correct formula.

The correct formula of a compound must represent its molecular weight. That is, the sum of the weights rep-

1 78 PRACTICAL CHEMISTRY

resented by the kind and number of atoms in the formula must be equal, or very nearly equal, to the molecular weight found by experiment. Let us take an example. A com- pound was found by analysis to contain 92.3 per cent of carbon and 7.7 of hydrogen, and to have a vapor density of 2.4375. Dividing the percentages by the atomic weights, we have: 92.3 -^ 12 = 7.7, and 7.7 -^ i = 7.7. Since 7.7:7.7 as 1:1, the compound contains at least one atom each of carbon and hydrogen, and its simplest formula is CH. This formula corresponds to the molec- ular weight 13. But the vapor density 2.4375 requires the molecular weight 78 {i.e. 2.4375 X 32), which is six times the weight (13) corresponding to the formula CH. Hence the molecular formula of this compound is not CH,

but CeHe.

We recapitulate regarding formulas of compounds : The simplest formula of a compound is found by dividing the per cent of each element by its atomic weight and reducing these quotients to the smallest whole numbers (if neces- sary). The molecular formula of a compound is found by three steps : (a) Find the simplest formula, (b) divide the molecular weight by the sum of the weights of the atoms in the simplest formula, (c) multiply the whole numbers of the simplest formula by the quotient obtained in (b).

If the molecular weight of a compound cannot be found by ex- periment, then the simplest formula is accepted as the molecular formula. For example, a compound contains 40 per cent of calcium, 1 2 of carbon, and 48 of oxygen. Dividing each per cent by the proper atomic weight, we have : 40 -i- 40 = i, 12 -7- 12 = i, 48 -^ 16 = 3. That is, one molecule of this compound contains (at least) one atom each of calcium and carbon, and three of oxygen; therefore the simplest formula is CaCOa- This is also accepted as its molecular formula, because the molecular weight cannot be found by any method known at present. (See last paragraph in 202.)

MOLECULAR WEIGHTS AND FORMULAS 179

207. Molecular formulas of elements. — The molec- ular weights of elements, as well as of compounds, can be found by experiment (202, 205). Such determinations show several interesting facts:

(i) Several gaseous elements have molecular weights which are twice the atomic weight. This means that the molecule consists of two atoms, and their molecular for- mulas are, for example, O2, Ho, CI2, N2 (also Br2).

(2) The molecular weights of many metallic elements and gaseous elements are identical with their atomic weights. This shows that the molecule and atom are iden- tical, and the molecular formula is the same as the atomic symbol, e.g. Na, K, Zn, Hg, Cd (cadmium), A (argon). He (helium), and Ne (neon). Such elements, especially the rare gases in the atmosphere, are often called monatomic gases to emphasize the fact that their molecules contain only one atom.

(3) Finally the molecular weights of certain elements vary with the temperature, decreasing with rise of tempera- ture, e.g. at lower temperatures, molecules of iodine, sul- phur, and phosphorus are represented by L, Sg, and P4, and at higher temperatures by I, S2, and P2.

208. Molecular equations. — Reactions involving the common elementary gases should be expressed by molec- ular equations. Thus, the equation for the formation of water vapor from hydrogen and oxygen is : —

2H2 + O2 = 2H2O

This equation is read thus : Two molecules of hydrogen unite with one molecule of oxygen to form two molecules of water vapor. Since this equation correctly represents the interacting substances as molecules, the equation is called a molecular equation. It should be noted that the

i8o PRACTICAL CHEMISTRY

proportions by weight are the same as in the simpler or gravimetric equation (135, 138).

Molecular equations are sometimes called volumetric equations or gas equations, because they show the volumes of gases involved in the reaction. Thus, the equation

H2 + CI2 = 2HCI may be written : —

H2 + CI2 = 2HCI

I vol. I vol. 2 vol.

because equal numbers of molecules represent equal vol- umes. This equation is read : One volume of hydrogen and one volume of chlorine form two volumes of hydro- gen chloride. It should be remembered that in molecular or volumetric equations a single molecule represents one volume of a gas (or vapor) and the coefficient indicates the number of volumes.

EXERCISES

1. Prepare a summary of molecules and molecular weights.

2. State and illustrate (a) Gay-Lussac's law and (b) Avogadro's theory.

3. (a) State the argument proving that a molecule of oxygen consists of at least two atoms. (&) Apply (a) to nitrogen.

4. Hydrogen and nitrogen combine in the ratio of 3 to i to form 2 volumes of ammonia. Show from this relation that a molecule of hydrogen contains at least two atoms.

5. What is the relation between molecular weight and vapor den- sity? Illustrate your answer.

6. Why is the formula of water vapor H2O and not HO or H2O2?

7. How are the molecular weights of gases determined by (a) the vapor density method, and (b) the mole method?

8. Define and illustrate (a) molecular equation, (b) mole, (c) gram- molecular weight, (d) gram-molecular volume.

9. What is a molecular formula? What is the molecular formula of oxygen, nitrogen, chlorine, hydrogen, zinc, mercury, sodium, potas- sium, argon, helium?

MOLECULAR WEIGHTS AND FORMULAS i8r

10. How is a molecular formula determined? Illustrate.

11. Express the following as volumetric equations : (a) One volume of phosphorus vapor and si.\ volumes of chlorine form four volumes of phosphorus trichloride (PClj) vapor; (b) carbon and water (vapor) form hydrogen and carbon monoxide.

12. Topics for review: (a) Gay-Lussac's law (93). (b) Equa- tions (Chapter X). (f) Atoms and molecules (101).

PROBLEMS

1. 1,000,000 molecules of hydrogen will unite with how many molecules of oxygen to form how many molecules of water vapor? What will be the relative weights of hydrogen and water vapor?

2. A liter of sulphurous oxide gas (SO2) weighs 2.9 gm. What is the molecular weight of this compound? Compare with the theoretical molecular weight.

3. The vapor density of hydrogen chloride is 1.14. Calculate the molecular weight.

4. If 3000 cc. of a carbon oxide weigh 3.75 gm., what is the molec- ular weight, formula, and name ?

5. Calculate the correct formula of the compound corresponding to (a) C = 39.Q, H = 6.7,0 = 53.4, vapor density = 1.906; (b) C = 73-8, H = 8.7, X = 17.1, vapor density = 5.03.

6. What volume of the constituent gases can be obtained by the complete decomposition of 6 1. of ammonia?

7. Write the equation for the reaction between nitric oxide and oxygen. What volume of oxygen is needed for 10 1. of nitric oxide? What will be the volume of the product?

8. What is the volume of (a) 2 moles of oxygen, (b) 0.5 mole of hydrogen chloride, (c) 22 gm. of carbon dioxide, {d) 22,400 cc. of hy- drogen?

9. The correct formula of a gas is CH4. (a) How many grams does a mole weigh? (6) What volume does this weight occupy? (c) What is the weight of i 1. of the gas?

10. A gas contains 69.49 per cent of oxygen and 30.51 of nitrogen, (a) What is its simplest formula? (b) If 500 cc. weigh 2.042 gm., what is the correct formula?

CHAPTER XVI ATOMIC WEIGHTS

209. Introduction. — In this chapter we shall study the methods of determining atomic weights. But first, let us review what we have already learned about atoms and atomic weights. In Chapter VII we learned that, according to Dalton's atomic theory (100), matter con- sists ultimately of particles called atoms, which have the same weight if alike but a different weight if different. These atoms, furthermore, remain undivided in chemical changes. Atoms unite and form molecules. The relative weights of atoms are called atomic weights. The standard for atomic weights is oxygen = i6.

210. Determination of approximate atomic weights from molecular weights. — The approximate atomic weight of an element can be determined from the molecular weights of several of its compounds.

We have seen (202, 205) that the molecular weight of a compound can be determined by experiment. After the molecular weights of several compounds have been de- termined, it is only necessary to find what parts of the molecular weights should be chosen as the atomic weight of the respective elements that constitute the compounds.

The steps in this method are : —

First, determine the molecular weights of several com- pounds of an element.

Second, find by analysis the per cent of the element in the compounds.

ATOMIC WEIGHTS

183

Third, find the weight of the element in each molecular weight by multiplying the molecular weight l)y the cor- responding per cent of the element.

Fourth, select the minimum value as the approximate atomic weight.

Let us find the atomic weight of carbon by this method. The steps are embodied in the accompanying table. Column I contains the name of each compound, column 2

Determixatiox of the Atomic Weight of Carbon

Per

MOL.

Cent

Wt. of

Moi..

Cent

Wt. OF

Compound

Wt.

Carbon

Compound

Wt.

OF CARBO^

Carbon

Carbon Monoxide

42.9

Ethylene

85.7

Carbon Dioxide

27.3

Acetylene

92.3

Methane

75-0

Ether

64-9

Ethane

80.0

Propane

81.8

the molecular weight, column 3 the per cent of carbon in each compound, and column 4 the weight of carbon in the corresponding molecular weight. The weights in column 4 are found by multiplying the molecular weight (in column 2) of the compound (in column i) by the corresponding per cent of carbon (in column 3) ; thus 28 Xo .429 = 12. The minimum value 12 is the atomic weight of carbon.

Obviously, the minimum weight must be the weight of a single atom, for it is highly probable that one or more compounds in a representative group will contain only one atom of a given element ; and the part of the molec- ular weight apportioned to this element will of course be its atomic weight. In the compounds that contain a multiple of this weight, it is likewise obvious that the

1 84

PRACTICAL CHEMISTRY

molecule must contain several atoms of the element. Thus, the weight of carbon in ethane is twice that in methane, and we conclude that a molecule of ethane contains two atoms of carbon — a conclusion in harmony with other observations.

The numerical results obtained by a study of the elements oxygen, hydrogen, chlorine, nitrogen, and carbon are summarized in the accompanying table. In this table, for the sake of simpHcity, whole numbers are used (except in the case of chlorine) and per cents are omitted. The minimum weight in each case is the atomic weight of the element, e.g. CI = 35.5.

Determixatiox of Approximate Atomic Weights

:-i z

•z,

H ^

^ ^

Compound

Molecular Weight

3 = 5

2^c

Weight of Chlorine

1^1

^ 0

* g

^ u

Water

■ —

—

Hydrogen Peroxide .

—

— ■

Hydrogen Chloride .

36.5

— ■

3S-5

Ammonia ....

I '7

—

Nitric Acid . .

—

Nitrous Oxide . .

—

Nitric Oxide . .

—

Nitrogen Dioxide .

—

Carbon Monoxide

—

Carbon Dioxide

—

Methane . . .

—

Ethylene . . .

—

Acetylene . . .

—

Ether ....

—

Ethyl Alcohol . .

—

Chloroform . .

II9-5

—

106.5

—

Carbon Tetrachloride

154

—

142

Cyanogen Chloride .

61.5

—

35.5

14 14

Minimum weight of each element .

35 5

1 12

This method of determining approximate atomic weights

ATOMIC WEIGHTS

was proposed about 1858 by the Italian chemist Canniz- zaro (Fig. Si).

211. Dulong and Petit's law. — Formerly the selection of the atomic weights of solid elements, especially metals, was checked by an approximate generaliza- tion commonly called the law of Dulong and Petit, which can be stated thus : —

The atomic weight is the quotient of 6.2 j {approxi- mately) divided by the specific heat.

As an example we take silver. Its specific heat is 0.057; if 6.25 is divided by this number, the quotient is about 109. This result shows that the atomic weight of silver is approximately 109 — not 55 or 21S. The exact value (found by other meth- ods) is 107.88.

More reliable methods are

now used in determining

. , ^ ,^1 1 ^, . FijT. 81. — Cannizzaro (1S26-1010)

atomic weights, although this *

so-called law is helpful in deciding between a number and its

multiple.

212. Equivalent weights. — The term equivalent urigJits rec^uires explanation. Equivalent weights of the elements are the weights that are chemically equivalent, provided the weights are expressed in terms of the standard for equiv- alent weights, viz. 8 grams of oxygen. In other words, the equivalent weight of an element is the weight that com- bines with or replaces 8 grams of oxygen. For example, magnesium and oxygen combine in the ratio of 1.5 to i respectively (in round numbers). Now if for the i we

1 86 PRACTICAL CHEMISTRY

substitute 8, then the 1.5 becomes 12 ; that is, 12 is the num- ber of grams of magnesium that unites with 8 grams of oxygen. Therefore 12 is the equivalent weight of mag- nesium. Similarly, hydrogen and oxygen combine in the ratio of i to 8 ; therefore i is the equivalent weight of hy- drogen.

The equivalent weights of certain elements are not found directly in terms of oxygen, but by finding the weight that combines with or replaces the equivalent weight of some other element. Thus, chlorine and hydrogen combine in the ratio of 35.5 to i ; therefore the equiv- alent weight of chlorine is 35.5. Sodium replaces hydrogen from water in the ratio of 23 to i ; therefore 23 is the equivalent weight of sodium. Again, 32.5 grams of zinc replace i gram of hydrogen from hydrochloric acid; therefore 32.5 is the equivalent weight of zinc.

These numbers are sometimes called combining num- bers, combining weights, or simply equivalents. The term equivalent weights is preferable, because they actually are the weights chemically equivalent to each other. Thus, if we start with hydrogen chloride (HCl), i gm. of hy- drogen — to take a convenient denomination — combines with 35.5 gm. of chlorine, and this i gm. of hydrogen can be replaced chemically by 32.5 gm. of zinc, 12 gm. of mag- nesium, 39 gm. of potassium, or 23 gm. of sodium, and so on. These elements are chemically equivalent in the ratio of these weights.

213. The relation of equivalent weights and atomic weights. — The equivalent weight of certain elements is equal to the atomic weight, e.g. hydrogen, chlorine, and sodium. In other cases the atomic weight is a multiple of the equivalent weight, e.g. 2 for oxygen, and 3 for alu- minium. We say, therefore, that the atomic weight of an element is equal to its equivalent weight or a small whole

ATOMIC WEIGHTS

187

multiple of it. This relation is seen in the accompanying table {iY\ which approximate weights are used).

Table of Equivalent Weights and Atomic Weights

Element

EQUrVALENT

Weight

Atomic Weight

Multiple

Oxygen

Aluminium

Bromine

Calcium

Carbon

Chlorine

Hydrogen

:Magnesium

Potassium

Silver

Sodium

Sulphur

Zinc

80 20

3 35-5

I 12

39 108

23 16

32.5

16 27 80 40 12

35-5

I 24

39 108

23 32 65

2 2

The equivalent weight of an element can be very ac- curately determined by experiment. Hence if we wish to determine the atomic weight accurately, we determine the equivalent weight and then multiply it by the correct whole number.

214. Accurate determination of atomic weights. — The method of determining atomic weights discussed in 210 in- volves finding the vapor density. But since the vapor density cannot be found accurately, atomic weights deter- mined by the vapor density method have only approximate values. The accurate value of an atomic weight is deter- mined by painstaking analysis of very carefully purified substances. The general method can be illustrated by an actual case kindly furnished by the American chemist Richards (Fig. 82), who has made exceptionally accurate determinations of atomic weights. In determining the

PRACTICAL CHEMISTRY

atomic weight of chlorine he found that 28.26299 gm. of silver chloride were formed from 21.27143 gm. of silver. He accepted AgCl as the formula of silver chloride and 107.880 as the atomic weight of silver, and calculated the atomic weight of chlorine thus : —

28.26299 — 21.27143 = 6.99156

\Vt. of : Wt. of : : At wt. of : At. wt. of

silver chlorine silver chlorine

21.27143 : 6.99156 : : 107.880 : x

X = 35-45S

The international atomic weight (35.46) is based on this and other determinations made by this same chemist.

215. International atomic weights. — An in- ternational committee se- lects the most accurate atomic weights of the elements. This commit- tee selected 16 as the atomic weight of oxygen (from several values formerly used). The weights are embodied in a table pubUshed at fre- quent intervals and called the International Table of Atomic Weights. The Fig. 82. — Richards (1868- ) ^^^j^ -^ g-^.^j^ ^^ ^}^g jn-

side of the back cover of this book. In this table the adopted atomic weights are placed in one column and the approximate weights in another. The approximate atomic weights are sufhciently accurate for general refer-

ATOMIC WEIGHTS 189

ences and in making chemical calculations ; they may be used in solving the problems in this book.

216. Review of methods of determining atomic weights. — The

methods and principles used in determining the atomic weight of an element can be reviewed by following the steps necessary in determin- ing the atomic weight of zinc (if its atomic weight were unknown).

(d) When zinc interacts with dilute hydrochloric acid, hydrogen is liberated ; and if a known weight of zinc is used, the weight of zinc needed to liberate i gm. of hydrogen is easily calculated. This num- ber is the equivalent weight of zinc. Now if one atom of zinc replaces one atom of hydrogen, the atomic weight of zinc and the atomic weight of hydrogen will have the same ratio as the weight of zinc and the weight of hydrogen found by experiment. , According to experi- ment the equivalent weight of zinc is about 32.5. This is its relation, atom for atom, to hydrogen, and thus far, is its atomic weight.

(6) When zinc and hydrochloric acid interact, zinc chloride is formed. If it is analyzed, the proportion of zinc to chlorine is about 32.5 to 35.5. Now if the elements combine, atom for atom, the atomic weight of zinc is 32.5 (assuming 35.5 as the atomic weight of chlorine) .

(c) However, when zinc is burned in air, zinc oxide is formed. If this compound is analyzed, the proportion of zinc to oxygen is about 65 to 16. If the elements combine, atom for atom, the atomic weight of zinc is about 65 (assuming 16 as the atomic weight of oxy- gen).

(d) According to these three determinations, the atomic weight of zinc is 32.5 or 65. We have assumed that the elements unite, atom for atom, in each compound. This is an incorrect assumption, be- cause an atom of zinc cannot have two different weights — 32.5 and 65. If the atomic weight is 32.5, zinc oxide must consist of one atom of oxygen and two of zinc. But if the atomic weight is 65, zinc chlo- ride must consist of two atoms of chlorine and one of zinc, and two atoms of hydrogen must have been replaced by one of zinc.

(e) The molecular weight of zinc* chloride is found by the vapor density method to be about 136. Zinc chloride contains 47.8 per cent of zinc. Therefore 47.8 per cent of 136, or 65.08, is zinc; that is, 65.08 is the approximate weight of the smallest part of zinc in zinc chloride, and it may be the approximate weight of one atom. If

iQO PRACTICAL CHEMISTRY

zinc chloride consists of two atoms of chlorine and one of zinc (weigh- ing 65), its molecular weight is about 136. In other words, it is evi- dent that our assumption regarding the number of atoms in zinc chlo- ride is probably correct.

(/) We are not absolutely positive, however, that the zinc in a molecule of zinc chloride may not be one atom weighing 65, or two atoms weighing 32.5 each. But the approximate atomic weight of zinc determined by applying the law of specific heats is 66.4 (i.e. 6.25 -^ 0.094). This shows clearly that the atomic weight of zinc is approximately 65 (and not 32.5).

(g) Accurate analyses of pure zinc compounds show that the atomic weight of zinc is 65.37.

EXERCISES

1. Prepare a summary showing the relation of atomic weights to (a) molecular weights and (b) equivalent weights.

2. State the three steps in finding atomic weights by the " mini- mum weight " method. Illustrate by carbon.

3. Describe Dulong and Petit's method of finding atomic weights.

4. What is an equivalent weight? Illustrate by aluminium.

5. State exactly the meaning of " chemically equivalent."

6. What is the equivalent weight of (a) oxygen, hydrogen, chlorine, carbon, sulphur, bromine, and (b) aluminium, calcium, magnesium, potassium, silver, sodium, zinc?

7. Describe the process of finding an exact atomic weight.

8. Why is 16 the atomic weight of oxygen?

9. What is the atomic weight of (a) Al, Ba, Br, Ca, C, CI, Cu, Y, Au, H ; [b) I, Fe, Pb, Mg, Mn, Hg, N, O, P, K ; (c) Si, Ag, Na, S, Sn, Zn?

PROBLEMS

1. If 2 gm. of potassium chloride yield 3.84 gm. of silver chloride, calculate the atomic weight of potassium.

2. If 91.46 gm. of metallic silver, when heated in chlorine, yield 121.4993 gn^- of pure silver chloride, calculate the atomic weight of chlorine.

3. Calculate the equivalent weights of the respective metals from the following data : (a) 0.5 gm. of calcium unites with 0.2 gm. of oxygen to form calcium oxide (CaO). (b) 15 gm. of mercury unite with 1.2 gm. of oxygen to form mercuric oxide (HgO).

4. Calculate the equivalent weight of sodium from the following :

ATOMIC WEIGHTS 191

(a) 2.3 gm. of sodium liberate o. i gm. of hydrogen from water.

(b) 1. 15 gm. of sodium liberate 555.5 cc. of hydrogen (at standard conditions).

5. If 0.03 gm. of magnesium yields 30.4 cc. of hydnjgen at 20" C. and 750 mm., what is the equivalent weight of magnesium?

6. Suppose 4.861 1 1 gm. of ferric o.xide (Fe-jOs) yield 3.39995 gm. of iron, what is the atomic weight of iron?

7. If 4.59507 gm. of phosphorus trichloride (PCI.3) give by pre- cipitation 14.381 18 gm. of silver chloride, what is the atomic weight of phosphorus?

8. In the synthesis of hydrogen bromide, 0.8606 gm. of hydrogen combined with 6S. 25033 gm. of bromine. Calculate the atomic weight of bromine.

9. By experiment, 14.36691 gm. of mercuric bromide (HgBr,) yield 14.9694 gm. of silver bromide. Calculate the atomic weight of mercury.

10. Assume that 1.70563 gm. of strontium bromide (SrBrj) require 1.48707 gm. of silver to precipitate the bromine as AgBr. What is the atomic weight of bromine ?

11. (a) What is the atomic weight of phosphorus, if the specific heat is 0.189? (b) Of potassium, if the specific heat is 0.166? (c) Of man- ganese, if the specific heat is 0.122? (d) Of tin, if the specific heat is 0.054? Compmre each with the exact atomic weight.

12. It was found that 4-58644 gm. of calcium bromide (CaBr>) require 4.95025 gm. of silver to precipitate the bromine as AgBr. What is the atomic weight of calcium?

CHAPTER XVII

VALENCE

217. What is valence ? — Formulas obtained by the meth- ods discussed in Chapter XVI show certain regularities. Let us take as examples the following groups : —

I. Hydrogen Compounds

HCl HoO

NH3

CH4

HBr HoS

PH3

SiH4

II. Oxi

ides

Na20

CaO AloOs

SO2

P2O5

SO3

K2O

MgO P2O3

CO2

AS0O5

Cr03

III. Acids and Salts

HCl

HNO3

H2SO4

H3PO4

NaCl

NaNOs

Na2S0,

Na3P04

CaClo

Ca(N03)2

CaS04

Ca3(P04)2

AICI3

A1(N03)3

AI2 (804)3

A1P04

IV. Bases

NaOH Ca(0H)2 A1(0H)3 KOH Ba(0H)2 Bi(0H)3

These groups might be greatly extended. Careful com- parison of these and many other formulas shows two sig- nificant facts, (i) Atoms of elements differ in the num- ber of atoms or atomic groups of the other elements with which they combine. Thus, one atom each of chlorine and bromine combines with one of hydrogen ; while one atom

192

VALENCE 193

each of oxygen and sulphur combines with two of hydrogen ; and so on. (2) One atom of certain elements unites with only one atom or atomic group of certain other elements, with only two atoms or two atomic groups of certain others, and so on. Thus, one atom of calcium combines with one atom of oxygen, two of chlorine, with two NO3 groups, one SO4 group, and two OH groups. Hence, we conclude that atoms of elements have definite combining capacity. The number which expresses the combining capacity of an atom of an element is called the valence of the element (Table A, 221).

218. Valence of radicals. — Strictly speaking, valence is a property of atoms. But many atomic groups often act chemically like individual atoms, e.g. they pass as a whole from one compound to another. Hence it is customary to assign a valence to atomic groups like SO4, NO3, NH4, and OH. Such groups are called radicals (163, 179) ; occasionally the term radical is applied by analogy to a single atom, e.g. CI in chlorides, C in carbides, S in sulphides (as in Table B, 221). In a word, a radical has a valence just as an element does.

219. Valence terms. — An element or radical which has the valence i is called a monad ; those w^hich have the valence 2 are called dyads; similarly, those elements or groups which have the valence 3, 4, 5, 6 are called re- spectively triads, tetrads, pentads, and hexads. The cor- responding adjectives are univalent, bivalent (or divalent), trivalent, quadrivalent (or tetravalent), quinquivalent (or pentavalent), and hexavalent.

220. How valence is represented. - The valence of an element or radical may be represented by writing a small Roman numeral slightly above the symbol, e.g. U\ O", Ar", SO4", OH'. Other ways are given in 226.

194

PRACTICAL CHEMISTRY

221. Tables of valence. — The valence of certain ele- ments and radicals is shown in Tables A and B. These tables are very useful in writing formulas (225).

Table A. — Valence of Certain Elements

Name

Symbol

Valence

Name

Symbol

Valence

Aluminium .

Iron (ous) . .

(Ammonium) .

(NH4)

Iron (ic) . . .

Antimony (ous)

Lead ....

Antimony (ic) .

Magnesium . .

Arsenic (ous) .

]\lercury (ous) .

Arsenic (ic) . .

Mercury (ic)

Barium . . .

Phosphorus (ous)

Bismuth . . .

Phosphorus (ic)

Calcium . . .

Potassium . .

Carbon . . .

Silicon . . .

Copper (ous) .

Silver ....

Copper (ic) . .

Sodium . . .

Gold (ous) . .

Tin (ous) . .

Gold (ic) . . .

Tin (ic) . . .

Hydrogen . .

Zinc ....

Table B. — Valence of Certain Radicals

Group Name

Symbol

Group Name

Symbol

COMPOLT«>

Formula

Compound

Formula

< >

Acetate ....

C2H3O2

Iodide

Bromide . . . .

Manganate . . .

Mn04

Carbide . . . .

Nitrate

NO3

Carbonate

CO3

Nitrite

NOo

Carbonate (acid) .

HCO3

Oxide

Chlorate . . . .

CIO3

Permanganate . .

Mn04

Chloride . . . .

Phosphate (ortho) .

PO4

Chromate . . .

Cr04

Silicate (meta) . .

Si03

Dichromate .

CrsO;

Sulphate ....

SO4

Ferricyanide . .

Fe(CN)6

Sulphate (acid) . .

HSO4

Ferrocvanide . .

Fe(CN)6

Sulphide ....

Fluoride . . . .

Sulphite ....

S03

Hydroxide . . .

Sulphite (acid) . .

HS03

VALENCE 195

For convenience in using or learning the usual valence of certain elements and the valence of radicals, the valences given in Tables A and B may be rearranged as follows : —

1. Monads (valence i) — Ag, Au (ous), Br (ide). CI (ide), CIO3, Cu (ous), F (ide), H, HCO3, HSO4, Hg (ous), I, K, MnO^ (per), Xa, NH4, NO2, NO3, OH.

2. Dyads (valence 2) — Ba, Ca, CO3, Cr04 ,Cr..07, Cu (ic), Fe (ous), Hg (ic), Mg, Mn04 (ate), O, Pb, S (ide), SO3, SO4, SiOs, Sn (ous), Zn.

3. Triads (valence 3) — Al, As (ous), Au (ic), Bi. Fe (ic), Fe(CN)6 (i). P (ous), PO4, Sb (ous).

4. Tetrads (valence 4) — C, C (ide), Fe(CN)6 (o). Si, Sn (ic).

5. Pentads (valence 5) — As (ic), P (ic), Sb (ic).

222. Variable valence. — Most elements and radicals have a fixed valence. But the valence of some elements varies with the combining element and the conditions under which the combination occurred. For example, iron, copper, mercury, and a few other ele- ments form two series of compounds in which these elements have a different valence. These series are designated as -ous and -ic (see 502). Thus, iron forms ierrous compounds {e.g. ferrous sulphate) in which iron has the valence 2, and ferr/c compounds {e.g. ferric chloride) in which iron has the valence 3. (See 228.)

223. Interpretation of valence by combination. — Ele- ments which have the same valence combine with each other unit for unit. This is a general rule. For example, sodium and chlorine each has the valence i ; hence one atom of sodium combines with one atom of chlorine to form one molecule of sodium chloride (NaCl).

Elements which have a different valence usually com- bine with each other so that the total valence of the atoms of each element in a molecule is equal. This is also a gen- eral rule. In other words, the valence of the two parts must balance. For example, two atoms of hydrogen (each having the valence i) combine with one of oxygen (having the valence 2) to form one molecule of water (HoO) ; sim- ilarly two atoms of aluminium (each having the valence 3)

196 PRACTICAL CHEMISTRY

combine with three atoms of oxygen (each having the va- lence 2) to form one molecule of aluminium oxide (AI2O3). The two rules just given apply to radicals. Thus, am- monium and hydroxyl each has the valence i ; so one NH4 radical combines with one OH radical to form one molecule of ammonium hydroxide (NH4OH). Similarly, two NH4 radicals combine with one S4O radical (having the valence 2) to form one molecule of ammonium sulphate ((NH4)2S04). So also, three atoms of calcium (each having the valence 2) combine with two PO4 radicals (having the valence 3) to form one molecule of calcium phosphate (Ca3(P04^'2)-

224. Interpretation of valence by displacement or sub- stitution. — ■ Just as atoms and radicals of the same valence combine, unit for unit, and those of different valence combine so that the valence of the two parts (element or radical) balances, so also atoms and radicals displace each other — unit for unit if the valence is the same, or equivalently if the valence is different. Let us take several examples : (i) When silver displaces hydrogen in hydrochloric acid, one atom of silver (having the valence i) displaces one atom of hydrogen (having the valence i), thereby forming silver chloride (AgCl). (2) Similarly, one atom of chlorine (having the valence i) displaces one OH radical (having the valence i) from sodium hydroxide, thus producing sodium chloride (NaCl). (3) So also, when aluminium is substituted for the hydrogen in sulphuric acid, two atoms of aluminium (each having the valence 3) are needed to balance the three SO4 radicals (each having the valence 2) in the aluminium sulphate (Alo(S04)3) formed.

225. Writing formulas from valence. — Formulas of many compounds may be written by using the tables of valence.

(i) Suppose we wish to write the formula of

VALENCE 197

J'rom 'ra])le A the valence of mag- nesium is found to be 2, and from Table B the valence of chlorine in chlorides is found to bei. Remembering the fnndamcntal rule that in most compounds the total valence of each of the two parts must be equal, it is clear that the formula must contain one atom of magnesium and two atoms of chlorine ; hence the formula is MgCb.

(2) Suppose we wish to know the formula of lead nitrate. From the tables, the valence of lead is 2 and of the NO3 radical is i ; so it is necessary to have two XO3 radicals for one atom of lead ; the formula, therefore, is

Pb(X03)2.

(3) Similarly, the formula of aluminium oxide is AI2O3, because 2 and 3 are the least number of atoms of Al and 0 which give a balanced valence (6 in each case).

Formulas of salts are frequently needed. These can be readily wTitten by learning the formulas of the common acids and substituting the correct number of atoms for the hydrogen. Thus, the formula of sulphuric acid is H0SO4. Therefore, all sulphates will have one or more SO4 radicals as one part of the formula and one or more atoms of a metal for the other part. If the metal has the valence of i, the formula will contain two atoms of the metal, e.g. Xa2S04, K2SO4. Ag2S04; if the valence is 2, there will be only one atom, e.g. CaS04, BaS04, CUSO4. Whereas if the valence is 3, there will be two atoms of the metal and three SO4 radicals, e.g. AI2 (504)3. Similarly, all nitrates have NO3 (one or more), all chlorides have CI (one or more), all phos- phates have PO4 (one or more), and so on. It is not neces- sary to learn by heart formulas of salts ; they can be written by applying the valence rules.

226. Representation of valence. — The valence of elements and radicals may be represented in various ways. One has already been

igS PRACTICAL CHEMISTRY

given, viz. H\ O", PO4'". Sometimes short lines are used, e.g. H — , 0 = , — O — , Al =, etc. If lines are used to represent valence in com- pounds, a single line answers for two elements. Thus, the formula

of water is written H — O — -H rather than H — — O H. Simi-

/H OH

larly, ammonia gas is N^H and calcium hydroxide is Ca<( or

\H \0H

,0— H Ca<^ . Such formulas are called structural or graphic formulas,

^O— H for they show the probable arrangement of the atoms in the mole- cules. Thus, the graphic formula of nitric acid is usually written

H— 0 — N.^' , because this arrangement of atoms not onlv shows

^O the correct valence of each element but it also represents certain facts about nitric acid, e.g. that the hydrogen is not combined di- rectly with nitrogen. Structural formulas are useful, especially in organic chemistry (Chapter XXIII), but it must not be forgotten that they are merely representations ; the lines are intended to in- dicate the numerical value of the valence and not the strength of the combination of the atoms.

227. Determination of valence from atomic and equiv- alent weights. — We learned in 213 that the atomic weight is equal to the equivalent weight or is a small whole multiple of it. The number (i or a small multiple) ex- pressing the relation of the atomic and equivalent weight of an element is the valence. Hence the valence of an element is found by dividing its atomic weight by its equivalent weight, that is: —

Valence = Atomic Weight -=- Equivalent Weight Thus, the atomic weight of magnesium is 24, and its equiv- alent weight is 12. Therefore the valence is 24 ^ 12, or 2.

228. Exceptional compounds. — The formulas of certain com- pounds apparently do not conform to the simplified conception of valence discussed above. Carbon monoxide (CO), calcium carbide

\aij:xck 199

(CaC-i), lead tctroxidc (Pb304), magnetic iron oxide (Fe304), hydrogen peroxide (H2O2), acetylene (C2H2),and ethylene (C2H4) are apparent exceptions. Interpretations of these apparent exceptions must be sought in a larger textbook. (See, however, 503, 573.)

229. Importance of valence. — Valence is very useful in writing formulas and in understanding analogous com- pounds. The valence of the common elements and radicals should be learned and used. (See especially Exercises 1, 3, 4, 8, 9, below.)

EXERCISES

1. What is the valence of these elements and radicals? (a) H, O, C, CI; {b) Ag, K, Na, NH4; (c) NO3, OH; (d) Ba, Ca, Cu, Fe (ous), Mg, Pb, Zn ; (e) CO3, SO4, S (ide) ; (/) Al, Fe (ic), PO4.

2. State the fundamental rule for writing formulas from valence.

3. Write the formula of the chloride of potassium, sodium, silver, copper (ous), copper (ic), mercury (ous), mercury (ic), iron (ous), iron (ic), zinc, tin (ous), tin (ic), calcium, barium, magnesium, bismuth, aluminium, ammonium, lead.

4. Write the formula of the sulphate of K, Xa, Ag, Cu, Fe (ous), Zn, Pb, Ca, Ba, Mg, Cr, Al, NH4.

5. Write the formula of the following compounds and indicate the valence of each element or radical by Roman numerals : ammonium fluoride, sodium silicate, potassium manganate, barium phosphate, zinc iodide, ammonium chromate, silver chromate, magnesium oxide, sodium dichromate, aluminium chloride, ferrous bromide, calcium phos- phate, mercurous nitrate.

6. Write the formulas of the following (as in Exercise 5) : ferrous carbonate, aluminium phosphate, calcium fluoride, sodium perman- ganate, phosphoric acid, silicic acid, sulphurous acid, nitrous acid, chromic acid, hydriodic acid, carbonic acid.

7. Define and illustrate the term eqiiivaleni weight. What is the equivalent weight of hydrogen, oxygen, sulphur, zinc, copper, magne- sium, silver, potassium, aluminium? State the relation of each to its atomic weight.

8. Write the formula of the nitrate of calcium, silver, iron (ic), mercury (ous), barium, magnesium, aluminium.

9. Write the formula of the oxide of Al, Sb (ous and ic). As (ous and ic), Ba, Bi, Ca.

200 PRACTICAL CHEAIISTRY

10. Write the formula of (a) the carbonate of Cu (ic), Ba, Ag;

(b) the iodide of Hg (ous and ic), Pb, Ag, Xa, Fe (ous and ic), Al ;

11. Write the formula of (a) the dichromate of Pb, Xa, Al, Ag ; (b) the manganate of Xa, Ca, Ag ; (c) the nitrate of Pb, Ca, Cu (ous and ic).

12. Write the formula of (a) the nitrite of Xa, Ca, Pb, Cu (ic) ; {b) the permanganate of Xa, Ag, H, K; (c) the sulphate, chloride, nitrate, carbonate, sulphite, bromide of H.

CHAPTER XVni IONS AND IONIZATION

230. Introduction. — Many properties of solutions and the general properties of acids, bases, and salts have al- ready been enumerated (161-165). Considerable atten- tion has also been given to examples of acids, bases, and salts, viz. hydrochloric and nitric acids, ammonium hydrox- ide, and chlorides and nitrates. In this chapter we shall consider more in detail the properties of dilute solutions of acids, bases, and salts.

231. Theory of ionization. — Solutions of acids, bases, and salts differ conspicuously in certain properties from solutions of other substances {e.g. sugar, glycerin, alcohol). That is, we have two classes of solutions. One class — acids, bases, and salts — is active chemically, conducts an electric current, and behaves abnormally when boiled or frozen. Whereas the other class — not acidic, basic, or salty — is not active chemically, does not conduct an elec- tric current, and behaves normally when boiled or frozen.

The properties of solutions, especially solutions of acids, bases, and salts, are explained more or less acceptably by the theory of ionization, which was proposed in 1887 ^Y the Swedish chemist Arrhenius (Fig. 83). We shall state and explain the theory first, and then show how the theory helps us understand the properties of dilute solutions of acids, bases, and salts. The theory is usually stated as follows : —

202

I'k \(l I( \L CIII'lMISrRY

Adtis, hdsrs. It 11(1 SiiHs, icl/cfi dissolved in Kuihr, d(r(f)n/>()sc

into /xirlidcs ( //tiri^nl u'/7// clcilriiily.

'This tlu'ory humus ihat ;i solution of sodium chloride,

lor cxaiupk', (.onsisls of water lhrouu;hout which arc dis- tributed souu' molecules of sodium ( hloridc to- gether with particles of electrically charged so- dium and chlorine into which the other mole- cules of sodium chloride ba\'e dissociated.

232. What are ions? - — The decomj)osition, or dissociation, of acids, bases, and salts when in solution is called ioniza- tion. The electrically charged particles formed bv ionization are called

Fig S^,. ArrluMuus (.850- ) .^^^g j^^^,j^ j^^,^ j^ .^ ^^^^^_

tion of a molecule. Two kinds of ions are present in every electrolytic solution, viz. electro-positive ions, or cations, and elect ro-negatixe ions, or anions.

K>ns, although formed by the dissociation of molecules, must not be confused with atoms. /^);/.s~ arc electrically charged atoms or radicals. The c>lectric charge changes the atom or radical, so that the properties of ions are quite tlitlerent from those of atoms. For example, in a solution of sodium chloride the electro-positive sodium ions move abcHit in the water witlunit producing any apjKirent chemical change ; whereas ordinary sodium interacts violently with water, as we have already seen (49). Similarly, the chloride

IONS AND IONIZATION 203

ions circulate freely in wdter and exhibit none of the effects of chlorine on water (145, 146 ^ It must be understood that the electric charges on the ions in a solution do not come from the electricity that may subsequently be passed through the solution. The ions are electrically charged just as soon as the molecules decompose in the solution (238).

233. How ions are represented. — Ions are represented by chemical symbols supplemented by the sign that desig- nates the kind and amount of the electric charge. Thus, the ions formed by sodium chloride are Na+ and Cl~. In sodium chloride solution a molecule of sodium chloride dis- sociates into one ion each of sodium and chlorine. Hence the number of sodium ions equals the number of chloride ions ; and the sum of the positive charges on all the sodium ions equals the sum of the negative charges on all the chlo- ride ions.

On the other hand in calcium chloride solution (CaClo), each molecule dissociates into two chloride ions and one calcium ion. Hence the ions formed by the dissociation of calcium chloride are designated Ca++ and 2C1~ {not Cl2~)*

Since the sum of each kind of electric charge must be equal {i.e. the total positive equals the total negative), the charge on each calcium ion must be twice the charge on each chloride ion. Similarly, the ions formed by cal- cium hydroxide (Ca(0H)2) are Ca++ and 2OH-, and by calcium sulphate are Ca^^ and S04~ ~.

In ordinary chemical formulas, atoms and radicals are represented as united, e.g. Cl-2, CaClo, Ca(0H)2. But when the molecule dissociates, independent particles (ions) are formed. Therefore we represent ions as separate particles. That is, the Clo in CaClo becomes 2CI- or Ch -h Ch, but not Cl2~.

234. A new definition of acid and base. — Specitic prop-

204 PRACTICAL CHEMISTRY

erties are exhibited by solutions of typical acids and bases. Thus, acids have a sour taste and turn Utmus red ; bases have a bitter taste and turn htmus blue. When these properties are interpreted by the theory of ionization, in- teresting and important conclusions result. Acids, bases, and salts are electrolytes. Hence their solutions contain ions, and the properties of such solutions are ascribed to the ions. According to the theory of ionization, then, an acid is a compound whose solution contains hydrogen ions (H+) , while a base is a compound whose solution contains hy- droxyl ions (0H-). These definitions should be compared with those previously given (161-165) .

235. Salts and ionization. — Salts may be defined in several ways. For example, they are compounds (other than water) resulting from neutralization of acids and bases. Again, they are formed by substituting (i) a metal for the hydrogen of an acid or (2) a non-metal for the hydroxy 1 of abase (163-165).

According to the theory of ionization, salts may be de- fined in two ways. First, salts are electrolytes which in solution yield neither hydrogen nor hydroxyl ions. Second, salts are compounds finally formed by the union of the posi- tive ion of a base and the negative ion of an acid. Later we shall study different kinds of salts. (See 273, 293.)

236. A new definition of neutralization. — In 164 we saw that neutralization is a chemical change in which an acid and a base interact and form a salt and water. Neutralization, interpreted by the theory of ionization, is the combining of hydrogen and hydroxyl ions to form molecules of water. Suppose solutions of hydrochloric acid and potassium hydroxide are mixed in the proper proportions. The mixture at first contains ions of hydro- gen, chlorine, potassium, and hydroxyl. But the hydrogen

IONS AND lONI/A'I ION 205

and hydroxyl ions immediately unite to form molecules of water, because water does not dissociate to any appre- ciable extent. The final solution is thus rendered neutral by the removal of the hydrogen and the hydroxy] ions — the acidic and basic constituents respectively.

The ionic equation expressing the neutralization of po- tassium hydroxide by hydrochloric acid is : —

K+ + OH- + H+ + CI- = K+ + CI- -f- HoO

The potassium and chloride ions move freely about in the solution. If the solution is evaporated, the ions unite as it becomes concentrated until nothing remains except the neutral salt potassium chloride. Since neutralization is the combining of hydrogen and hydroxyl ions to form water, the gen^^ral ionic equation for neutralization is : —

H+ -f OH- = H,0

Hj'drogen Ion Hydroxyl Ion Water

237. Interpretation of certain facts by the theory of ionization. — ■ The theory of ionization, like other theories, must meet one import

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