IRLF B 3 TDM CHEMISTRY AND ITS RELATIONS TO DAILY LIFE THE MACMILLAN COMPANY NEW YORK BOSTON CHICAGO DALLAS ATLANTA SAN FRANCISCO MACMILLAN & CO., LIMITED LONDON BOMBAY CALCUTTA MELBOURNE THE MACMILLAN CO. OF CANADA, LTD. TORONTO JUSTUS VON LIEBIG. 1803-1873. The father of Agricultural Chemistry. He was the first to teach rational basis for fertilizing the land. CHEMISTRY AND ITS RELATIONS TO DAILY LIFE A TEXTBOOK FOR STUDENTS OF AGRICULTURE AND HOME ECONOMICS IN SECONDARY SCHOOLS BY LOUIS KAHLENBERG PROFESSOR OF CHEMISTRY AND DIRECTOR OF THE COURSE IN CHEMISTRY IN THE UNIVERSITY OF WISCONSIN AND EDWIN B. HART PROFESSOR OF AGRICULTURAL CHEMISTRY AND CHEMIST TO THE AGRICULTURAL EXPERIMENT STATION IN THE UNIVERSITY OF WISCONSIN gorfe THE MACMILLAN COMPANY 1913 All rights reserved COPYRIGHT, 1913, BY THE MACMILLAN COMPANY. Set up and electrotyped. Published June, 1913. Noroiooti J. S. Gushing Co. Berwick & Smith Co. Norwood, -Mass., U.S.A. PREFACE THIS book is intended to represent a year's work for students of agriculture and home economics in secondary schools. The aim has been to make the subject matter thoroughly practical in character and to present it in an interesting and simple way so that the student may grasp it. At the same time, the delineation has been made on sound scientific lines, and it will require patient and con- tinuous application on the part of the student to accomplish the work intelligently. Useful facts have naturally been placed in the foreground, and no more theory has been presented than necessary. Chemical formulas have been in- troduced to some extent, but merely as an aid in expressing facts in simple, compact, and convenient form. Atomic and molecular theories have not been presented, for they are not calculated to aid students at this stage of advancement. While the book is not intended as a preparatory course for colleges, yet those who have completed it successfully will doubtless have gained information and sound scientific training which will compare favorably with that received in the pursuit of courses usually taken in preparatory schools. The laboratory experiments detailed in Chapter XXI are to be performed by the student in connection with the study of each of the preceding chapters. The teacher should super- vise this work carefully and require the student to keep neat, accurate records in a notebook. At the end of Chapter XXII are given lists of apparatus and chemicals that will be needed. Addresses of a few well-known firms from whom V 265529 vi PREFACE such supplies may be obtained have also been furnished for the convenience of the teacher. The questions at the end of each chapter will help in reviewing the salient points of that chapter. The teacher will, of course, raise many additional questions, especially in connection with the laboratory exercises, upon which special stress should be laid. While this book is primarily intended as a textbook to be used by pupils in schools under the supervision of a teacher, there are doubtless many others who will find it helpful as a volume of general information for home reading and study. Any corrections or suggestions that may be use- ful in preparing future editions will be gratefully received. THE AUTHORS. MADISON, WISCONSIN, February 20, 1913. CONTENTS CHAPTER I. II. III. IV. V. VI. VII. VIII. IX. X. XI. XII. XIII. XIV. XV. XVI. XVII. XVIII. XIX. XX. XXI. XXII. INDEX GENERAL FUNDAMENTAL CONSIDERATIONS THE COMPOSITION AND USES OF WATER . HYDROGEN, OXYGEN, HYDROGEN PEROXIDE, AND OZONE THE AIR, NITROGEN, NITRIC ACID, AND AMMONIA ACIDS, SALTS, BASES, AND CHEMICAL FORMULAS . THE HALOGENS SULPHUR, PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH BORON AND SILICON . CARBON AND ITS COMPOUNDS THE METALS OF THE ALKALIES AND THE ALKALINE EARTHS ALUMINUM, THE HEAVY METALS, AND THEIR IM- PORTANT ALLOYS PAINTS, OILS, AND VARNISHES ..... LEATHER, SILK, WOOL, COTTON, AND RUBBER THE SOIL COMMERCIAL FERTILIZERS FARM MANURE PLANT LIFE AND WHAT IT PRODUCES THE ANIMAL AND ITS FEED HUMAN AND ANIMAL FOODS ..... MILK AND ITS PRODUCTS POISONS FOR FARM AND ORCHARD PESTS PRACTICAL LABORATORY EXPERIMENTS PAGE 1 22 30 41 52 59 74 81 118 140 171 179 191 214 229 238 255 269 291 307 321 381 vii CHEMISTRY AND ITS RELATIONS TO DAILY LIFE i CHAPTER I GENERAL FUNDAMENTAL CONSIDERATIONS ANIMALS, plants, water, rocks, soil, and the air are the things that make up the earth and its atmosphere. These have been the subject of study since earliest times, and a considerable amount of knowledge has consequently been gathered concerning living beings, plants, and animals, and their relations to water, the air, and the solid portions of the earth's crust. All of this knowledge is the result of observa- tion by means of our senses, sight, hearing, smell, taste, touch, and experience of changes of temperature. For many cen- turies the results of such observations have gradually been collected, corrected, and recorded in books and manuscripts. Such carefully gathered and systematically arranged knowl- edge is called Science. In other words, science is systematized knowledge resulting from careful observation made with a definite purpose in view. The opinions that men have formed as a result of reflection, that is, thinking over such observa- tions, also form a part of science, and indeed such opinions are commonly recorded in scientific books. Now it is clear that these opinions may differ somewhat, and sometimes they differ quite widely, especially if the observations have been made with insufficient care, or if not enough observations have been made to warrant drawing a fairly safe conclusion. 2 CHEMISTRY AND DAILY LIFE The opinions then change from time to time as better obser- vations and more of them accumulate. Consequently the opinions, that is to say, the theories or theoretical views, are the changeable, the ephemeral part of science, though, to be sure, the observations too may be improved and multi- plied, as already stated. The Greek philosopher Aristotle (384-322 B.C.) taught that water, air, earth, and fire are the elements of which all other things are composed, and indeed this view was held for many centuries. Even now in literature, in the descriptions of storms, conflagrations, and the like, we sometimes still speak of the " battle of the elements." Such allusions, to be sure, are now only figures of speech, for the elements of the ancient Greeks are not at all the elementary constituents of which things are composed. In fact a careful study of all of the available materials on the earth, including the atmosphere has shown that there are about eighty substances that cannot be made from other substances and cannot themselves be resolved into anything else. For example, iron is one of these eighty substances. No one has ever been able to make iron from other substances that do not already contain it; neither has it been possible to produce other substances from iron alone. Iron is consequently spoken of as an element or an elementary substance. Similarly carbon is an element. This substance we have in almost pure form in the diamond, graphite, and the better grades of charcoal. In Table 1 given on the opposite page is a list of the chemical elements and the abbreviations or symbols by means of which they are very often designated. In the list there are many commonly known substances. These have been designated by means of heavy type. These more common elements number only about twenty- five, and indeed the most important elements are scarcely GENERAL FUNDAMENTAL CONSIDERATIONS TABLE 1 THE CHEMICAL ELEMENTS AND THEIR SYMBOLS ELEMENT SYMBOL ELEMENT SYMBOL Aluminum Al Neodymium Nd Antimony Sb Neon Ne Argon A Nickel Ni Arsenic As Niton (radium Barium Ba emanation) Nt Bismuth Bi Nitrogen N Boron B Osmium Os Bromine Br Oxygen Cadmium Cd Palladium Pd Caesium Cs Phosphorus P Calcium Ca Platinum Pt Carbon C Potassium K Cerium Ce Praseodymium Pr Chlorine Cl Radium Ra Chromium Cr Rhodium Rh Cobalt Co Rubidium Rb Columbium Cb Ruthenium Ru Copper Cu Samarium Sa Dysprosium Dy Scandium Sc Erbium Er Selenium Se Europium Eu Silicon Si Fluorine F Silver Ag Gadolinium Gd Sodium Na Gallium Ga Strontium Sr Germanium Ge Sulphur S Glucinum Gl Tantalum Ta Gold Au Tellurium Te Helium He Terbium Tb Hydrogen H Thallium Tl Indium In Thorium Th Iodine I Thulium Tm Iridium Ir Tin Sn Iron Fe Titanium Ti Krypton Kr Tungsten W Lanthanum La Uranium U Lead Pb Vanadium V Lithium Li Xenon Xe Lutecium Lu Ytterbium Magnesium Mg (Neoytterbium) Yb Manganese Mn Yttrium Yt Mercury Hg Zinc Zn Molybdenum Mo Zirconium Zr 4 CHEMISTRY AND DAILY LIFE more than a dozen. All known substances, natural or arti- ficial, whether they exist in the atmosphere, as part of the earth's crust, in natural waters, as parts of plants or animals, or indeed as parts of the heavenly bodies, are composed of one or more of the chemical elements mentioned in Table 1. It is the task of the chemist to study the composition of sub- stances, to build up new substances from given ones, and again to tear down complex substances, thus resolving them into simpler bodies. Whenever a change takes place in which new substances are formed, the change is called a chemical change. So, for example, when a piece of paper is burned, entirely new substances are formed ; again, when iron rusts, a new substance, the rust, is produced. These are examples of chemical change. The chemist, however, studies not only the new products that are produced from other substances, but he also takes into consideration the means employed to cause the chemical change to take place. Then again he studies the rate with which the change proceeds, ascertaining also whether all or only part of the substance has finally been transformed. The various conditions that affect the rate with which the change takes place, and the accompaniments of the change such as alterations of temperature, volume, evolution of light, electricity, etc., are also carefully studied. The sum total of all these observations, systematically arranged, constitutes the science of chemistry. It touches every walk of life, for wherever there is anything material, that is to say, wherever there is anything that we can touch, hear, see, taste, or smell, we have an object with which chemistry is con- cerned. In fact chemical changes are continually going on about us. The food we eat is chemically transformed into the various tissues of our bodies, and these in turn are acted upon by oxygen derived from the air we breathe, and thus oxidized into simpler substances which are in turn eliminated GENERAL FUNDAMENTAL CONSIDERATIONS 5 by means of the breath, the bowels, the kidneys, and the skin. Thus our very lives depend upon a series of chemical changes. Similarly the lives of all animals and plants depend upon chemical transformations of the material they take into their bodies, the products of such chemical changes being finally eliminated as the organism lives. But chemical changes are continually in progress in inanimate things, also. Thus the soil and rocks are being acted upon continually by water and the gases of the atmosphere, new products result- ing, which are more or less soluble in the water. Materials from these solutions are taken up by the rootlets of trees and other plants. Furthermore as natural waters are drunk by animals, some of these soluble substances are absorbed by the membranes of the digestive tracts of the organisms and used in building up new tissues. Many industrial processes and common operations in the household and on the farm depend upon chemical changes, and it is the purpose of this book to study some of the more important of these. It should be stated here, however, that whenever a change takes place which is not accompanied by a formation of one or more new substances, the change is called a physical change. Thus when a piece of gold is heated, it still remains gold and there is no change except a change of tem- perature and of volume. Again, when a glass rod is electrified by rubbing it with a silk rag, neither the substance of the rag nor that of the glass is altered. These, then, are typical examples of physical change. ' Physical changes are exceed- ingly numerous, and they very frequently occur without any accompanying chemical change. On the other hand, chem- ical changes practically never occur without accompanying phys- ical changes, for the formation of new substances is generally accompanied by increase or decrease of volume, rise or fall of temperature, and frequently by evolution of light, electricity, 6 CHEMISTRY AND DAILY LIFE or still other forms of energy. So when a piece of magnesium is burned in the air, a white powder, magnesium oxide, is formed, and this is more bulky than the original magnesium ; furthermore, during the burning of the metal both light and heat are evolved. A glance at Table 1 shows that the elements may be roughly divided into two great classes, the metals and the non-metals. It is, however, not possible to make a perfectly sharp division into these two classes, for there are some elements that exhibit properties which partake of the nature of both the metal and the non-metal. Thus it would be easy to class such elements as copper, silver, iron, gold, mercury, zinc, and sodium as typical metals ; and again, it would be simple to place such elements as sulphur, oxygen, carbon, chlorine, and phosphorus as non-metals. But elements like arsenic and antimony, for example, while possessing metallic luster and also some of the chemical properties of typical metals, are nevertheless brittle and in their chemical behavior also otherwise closely related to the non-metals. Such elements as partake of the nature of both the metals and non-metals are sometimes called metalloids. They are in reality a transition between the metals and the non-metals. When a substance consists of two or more elements, it is called a compound. Thus oxide of magnesium is a compound, for it is formed by the union of oxygen and magnesium. Lime is a compound, for it contains the metal calcium and oxygen. Sugar is a compound, for it consists of oxygen, hydrogen, and carbon. Water is a compound, for it is composed of oyxgen and hydrogen. As water is the most common compound on the earth, and as the elements that compose it are also of vast importance in all plant and animal life, the chemistry of water 'will be our first object of study. GENERAL FUNDAMENTAL CONSIDERATIONS 7 QUESTIONS 1. What is science? 2. What is a fact? 3. What is a theory? t 4. Mention the things which the ancient Greeks called elements. 5. What is a chemical element as regarded at present? 6. About how many chemical elements are there, and how may they be classified? 7. What is a compound? Give three examples. 8. What is a physical change? Give three illustrations of physical changes. 9. What is a chemical change ? Give four examples of chemical change. 10. Do chemical changes occur with or without accompanying physical changes? Explain your answer by means of a specific illustration. FIG. 1. CHAPTER II THE COMPOSITION AND USES OF WATER WATER is a compound of the ele- ments oxygen and hydrogen. When an electric current is passed through water in an apparatus like that shown in Fig. 1, oxygen is evolved on the positive plate and hydrogen on the negative plate ; and it is found that the volume of the oxygen is one half that of the hydrogen. That is to say, when water is decomposed by electrolysis, one volume of oxygen and two volumes of hydrogen are formed. Now one volume of oxy- gen weighs very nearly sixteen times as much as an equal volume of hydro- gen; consequently the vol- ume of oxygen produced in the electrolysis of water weighs eight times as much as the two volumes of hy- drogen that are liberated. By taking two volumes of hydrogen and one vol- ume of oxygen, mixing them, and igniting the mix- The electrolysis of water. ture with a flame or an elec- THE COMPOSITION AND USES OF WATER 9 trie spark, as, for example, in the apparatus in Fig. 2, water is formed and no hydrogen or oxygen is left uncombined. Thus the qualitative and quantitative composition of water is clearly established. FIG. 2. Explosion of a mixture of hydrogen and oxygen. Though water is extremely abundant, it never is found in the pure state in nature. The waters of the oceans and many inland seas are salty. Lake water and the waters of brooks and rivers, as well as those of springs and wells, always contain more or less solid saline material which has been dissolved by the water while in contact with the soil and rocks, On evaporating such saline waters, all of the solids remain behind as a residue. Thus when sea water or well water is placed in a dish and allowed to evaporate either slowly at 10 CHEMISTRY AND DAILY LIFE ordinary temperature or upon boiling, there remains a residue of the material that was dissolved in the water. By evaporating several gallons of water one may obtain a suffi- cient amount of such dissolved matter to make a chemical analysis of it and thus determine its exact nature. Thus the dissolved matters in many terrestrial waters have been analyzed. Both the nature and the amounts of the saline mat- ters contained in a water are determined by the character of the rocks over which the water has coursed. The purest natural water is rain water. It contains no mineral matter, except such small quantities of dust as have been carried into the air by the wind and then washed down by the descending drops. After it has rained awhile and the air has been well washed, the rain that falls consists of much purer water, but it always contains air dissolved in it and not infrequently nitrogenous compounds like nitrites and nitrates of ammonium, which are formed as the lightning flashes through the air. By boil- ing water and condensing the vapors a fairly pure water, com- monly called distilled water, is obtained. Figure 3 shows a simple apparatus for producing distilled water. Every distill- ing apparatus consists of a retort, a condenser, and a receiver. In Fig. 3 these are indicated by A, B, and C, respectively. The heat of the sun causes water to evaporate continually from the surface of the oceans, lakes, rivers, and moist soil. This moisture is then condensed on reaching cooler layers of air, forming clouds, fogs, dew, and rain. The rain water permeates the soil, dissolves portions of it, and gradually makes its way into brooks, rivers, lakes, and the sea. Then evaporation again goes on from the surfaces of these bodies of water, condensation again takes place, rain falls, and so the ceaseless round continues as long as the sun furnishes the needed energy. Thus it is clear that all water power is really derived from the sun. But not all of the water is thus THE COMPOSITION AND USES OF WATER 11 evaporated or carried into the sea. Some of it, though, to be sure, but a relatively small portion, is taken from the soil by the rootlets of plants, and again a certain amount is drunk by animals. Without water, plants and animals cannot live. Indeed the bodies of all plants and animals contain from fifty to ninety per cent of water by weight. This water, to be sure, is more or less firmly combined with the other materials that en- ter into the composition of the plant and animal bodies. From FIG. 3. Tho distillation of water. the latter it may be obtained by desiccation. That the tis- sues of plants and animals, soils, and not infrequently min- erals, rocks, and various artificial salts like blue vitriol, epsom salts, etc., contain water may readily be shown by gently heating a small portion of the substance, about the size of a small nut, in a test tube as shown in Fig. 4. The water set free condenses in drops on the upper cooler portions of the test tube. In such experiments one soon discovers that some compounds retain the water much more tenaciously than others, for while very slight warming suffices to liberate the water in some instances, other substances need to be heated even to redness before the water is set free. So, for 12 CHEMISTRY AND DAILY LIFE example, when grass, a piece of carrot, or lean meat is gently heated in a test tube, water is readily formed in drops on the upper cooler parts of the tube. To drive the water out of an old piece of mortar or con- crete would, however, require a much higher temperature. In- deed, it would be necessary to heat the substance to dull red- ness. The amount of water contained in plants and animals is usually high. But different organisms frequently contain quite different FIG. 4. -Testing^ substance quantities of water, the amount of which is commonly found by weighing the plant or animal substance and then drying it to constant weight at the boiling point of water, namely, 100 C. The difference between the original weight and that of the dried residue represents the moisture content. In this way it has been found that the leaves of herbaceous plants contain from 60 to 80 per cent water; potatoes, juicy fruits, and succulent plants contain from 85 to 95 per cent water; and algae and other aquatic plants may even contain as much as 98 per cent. On the other hand, wood commonly contains only from 44 to 55 per cent water, and many dry seeds contain much less. Dry grass seed, for ex- ample, contains only about 15 per cent water. The dried residue of the plant substance may be burned, whereupon there remains the ash, which represents the mineral matter in the plant. The combustible portion, which has been de- stroyed, is the so-called organic matter of the plant. It is oxidized and largely volatilized in the process of burning THE COMPOSITION AND USES OF WATER 13 (see combustion), there being formed gaseous substances like carbon dioxide, water vapor, nitrogen, and ammonia. The amount of ash found in different plant substances also varies considerably. Computed on the dried material, for example, oak wood contains about 0.48 per cent ash, rye straw 4.46 per cent, rye grain 2.09 per cent, tobacco leaves 17.16 per cent, potato leaves 8.58 per cent, potatoes 3.79 per cent. In general the leaves of plants contain more ash than the other parts. The ash of different plants has different composition, which fact will be discussed more fully later. The water content of animals also varies considerably. The following table indicates the amount of water contained in a few of the common animals, the entire body being considered : TART F 2 ANIMAL WATER CONTENT Fat ox 45.5 per cent Fat calf 63 per cent Fat pig 41 per cent Lean pig 55 per cent Fat sheep 43 per cent Lean sheep 51 per cent Chicken flesh 74 per cent Goose flesh .42 per cent Turkey flesh 55 per cent Brook trout 78 per cent White fish 70 per cent Lobster (exclusive of shell) 79 per cent Oyster (exclusive of shell) 81 per cent Table 3 presents the water content of various parts of an ox. PART WATER CONTENT Brain 80 per cent Heart 63 per cent Muscle of the loin . ' 62 per cent Kidney ' 76 per cent Liver 71 per cent Sweet breads 71 per cent Lungs 80 per cent 14 CHEMISTRY AND DAILY LIFE From these data it is apparent that considerably over half of the weight of animal tissues consists of water. It should be borne in 'mind, however, that this water is more or less tightly combined with the other constituents of the tissues. By expelling the water in the process of drying, these combi- nations are ruptured, the water being thus liberated. Now as human beings subsist on food that comes from animals and plants, it is clear that the human body derives from such food a considerable amount of the water it needs. This, however, is not sufficient, for drinking water is required in addition. Chemically pure water is not good drinking water. Freshly distilled water, obtained by careful distillation of spring or artesian well water, is insipid and flat to the taste, chiefly because by the process of boiling during distillation the air which was dissolved in the water has been expelled. For the same reason water that has been boiled and then cooled without sufficient opportunity to become thoroughly aerated again has a flat taste. Many natural spring waters and well waters are excellent drinking waters, though they always con- tain solid matter in solution that comes from the rocks and soils over which the water has flowed. The amount of this mineral matter varies very greatly with the nature of the solid material with which it has been in contact. So, for instance, water that has coursed over granite rocks only has but very little mineral matter in it, for granite is but slightly soluble in water. On the other hand, water in limestone regions is highly charged with mineral matter, for limestone is relatively readily soluble in water especially when the latter contains carbon dioxide, and this is always the case with water that has been in contact with the air, for the latter contains carbon dioxide (see air). Table 4 gives the amount of mineral matter contained in a few typical natural waters in grams per 1000 liters. These THE COMPOSITION AND USES OF WATER 15 figures are obtained by evaporating a given weight of water to dryness, weighing the residue, and then computing the amount of the latter per 1000 liters of the water taken. TABLE 4 MINERAL CONTENT IN WATER FROM GRAMS PER 1000 LITERS Atlantic Ocean 35,664 Indian Ocean 35,525 White Sea 33,118 Dead Sea 253,016 Great Salt Lake 302,122 Lake Michigan 145 Vichy Springs 6,798 Karlsbad Springs 6,091 Artesian well at St. Petersburg 3,891 Artesian well at London 834 Artesian well at Madison, Wis 371 Spring near Wausau, Wis 64 Rhine River 231 Nile River . . ' 142 It is to be borne in mind that the mineral matter dissolved in river waters often varies materially according to the time and place of taking the sample. The mineral content of waters of lakes and oceans also varies somewhat with the depth and locality. The figures in Table 4 are consequently to be regarded as subject to certain* fluctuations. They serve, however, to give a general idea of the amount of mineral matter found in natural waters. The mineral content of oceanic water and of salt lakes and seas consists chiefly of common salt, and such water is, of course, not fit to drink. Spring and well waters that contain considerable amounts of calcium salts (also often called lime salts) in solution are called hard waters. They may be quite good for drinking purposes, but they are not good for washing or steam boiler purposes, for in the boilers they deposit a hard coating or scale, like that commonly found in 16 CHEMISTRY AND DAILY LIFE an ordinary teakettle, for example, and with soap they form no suds, but simply a curd, which is in reality an insoluble calcium soap resulting from the decomposition of the soluble soap employed, by the calcium salts in the water. Before such water can be used for washing or boiler purposes, the lime salts it contains must be removed. This can be done in one of two ways : (1) by distilling the water, a process which on account of the fuel required is expensive and so practically out of the question for ordinary purposes, and (2) by throwing the lime salts out of solution by adding some other ingredients which will form insoluble compounds with the lime in the water. Soluble salts which thus act on the lime salts and form precipitates with- them are often spoken of as water cleansers or water purifiers. Among those commonly in use are sodium carbonate (also called washing soda), borax, and sodium phosphate. Of all of the natural waters rain water is by far the best and cheapest for washing and boiler purposes. It should be collected and stored in clean tanks or cisterns to which air has ready access. By the use of Portland cement (which see) it is now an easy matter to construct suitable reservoirs for the storage of rain water at a moderate cost. Such cistern water when clean and well aerated is good drinking water, though its mineral content is quite low, being simply due to the small amount of material dissolved from the cement-lined walls of the cistern. Sometimes spring water of well water is found to be unfit to drink because of the poisonous or otherwise injurious mineral ingredients it contains. But such cases are after all quite rare, being in general confined to excessively alkaline or strongly saline waters. For example, near copper deposits the waters may contain a considerable amount of compounds of that metal and so be injurious to health. In any case ic hen water has a taste or smell that causes suspicion, it ought THE COMPOSITION AND USES OF WATER 17 to be properly tested to ascertain whether it is potable. Water is more frequently rendered unfit to drink because of contami- nation with plant or animal refuse, sewage, disease-producing FIG. 5 (a). The kind of a well into which surface water will readily seep. FIG. 5 (6). An example of how pollution may occur. bacteria, or other pathogenic organisms. Wherever there is decaying animal or vegetable matter there microorganisms abound, and water contaminated with such material is dangerous 18 CHEMISTRY AND DAILY LIFE to health. A drinking water ought consequently to be ex- amined as to its content of microorganisms as well as sub- jected to chemical tests before being pronounced safe for drinking purposes. Surface wells near human habitations soon become contaminated from seepage from outhouses, cesspools, barnyards, and the like. It is to be borne in mind that after all a well is merely a hole in the ground into which water from the surroundings tends to drain. Mounding up the earth around the pump or opening of the well, of course, does not prevent the water of the soil in the neigh- borhood from getting into the well. Such a mound at the opening of the well can at best only prevent the water that is on the surface of the ground from running into the mouth of the well. Typhoid fever is one of the commonest and most dan- gerous diseases that may re- sult from drinking polluted water. This disease is due to the bacillus typhosus, which is not infrequently present in contaminated water. Dysentery and other intestinal and gastric disturbances may result from the presence of other organisms. // there is no alternative but to use water that is known to be contaminated for drinking purposes, it should first be boiled. This destroys the organisms it contains, but the water after boiling tastes flat as already mentioned. Con- taminated water is also frequently filtered through unglazed porcelain. Such filters are commonly known as the Pasteur FIG. 6. Typhoid bacillus greatly magnified. THE COMPOSITION AND USES OF WATER 19 water filters. They remove nearly all of the matter sus- pended in the water as well as a very large share of the organisms that are present. But such filters need to be renewed frequently, for the refuse they accumulate in their pores soon becomes only an added source of danger toward further con- tamination. On a large scale water is frequently purified by filtration through prop- erly constructed filters of gravel or crushed rock and sand. These too are fairly efficient in removing danger- ous material suspended in the water, but they also require to be renewed from time to time. Water derived from wells that are a hun- dred feet or more deep has obviously been well filtered in coursing through the vari- ous strata to the depth from which it issues, and conse- quently such water is practi- cally always free from organic contaminations. Deep well water is therefore usually to be preferred for drinking purposes. Besides resorting to boiling or filtration for the purpose of FIG. 7. Pasteur water filter. 20 CHEMISTRY AND DAILY LIFE removing organic contamination from a potable water supply, chemical means of purification may be adopted. This always involves introducing some chemical substance into the water, which destroys the organisms it contains. For such purposes a solution of sodium or calcium hypo- chlorite (bleaching powder) is not infrequently employed, especially in the case of a city water supply where the water in the reservoir and mains shows a high bacterial content. It is best to avoid the introduction of such chemicals into drinking water, for these substances which destroy bacteria are also deleterious to health. It is only when the disease germs are actually present in the water that such chemical treatment is really justifiable, for the germs are far more dangerous than the relatively small amounts of the anti- septics required to kill them. An antiseptic is a chemical that destroys germs, or prevents or retards their growth, without exerting a materially harmful effect upon the living body. Antiseptics therefore frequently are germicides. These will be referred to again. QUESTIONS 1. How may it be shown that water is a compound of oxygen and hydrogen ? 2. In what proportions by weight do these elements occur in water ? In what proportion by volume ? 3. Why are natural waters not chemically pure? What im- purities are there in rain water ? In well water ? 4. Draw a diagram of an apparatus for making distilled water, naming the principal parts of the apparatus. 5. Make a list of the important ways in which water is found in nature. 6. How may one determine whether a substance contains water ? Give a specific example. THE COMPOSITION AND USES OF WATER 21 7. About how much water is there in an animal like a cat ? In a dahlia root ? 8. What are some of the important characteristics of a good drinking water ? 9. What is meant by the term hard water ? 10. How would you treat a water that is known to be polluted and yet must be used for drinking purposes ? Why ? CHAPTER III HYDROGEN, OXYGEN, HYDROGEN PEROXIDE, AND OZONE WHILE both hydrogen and oxygen are obtained when the electric current is passed through water, these gases may also be prepared by other methods. On account of the frequent occurrence of these elements in many common substances, it is necessary to study their properties somewhat more closely. Hydrogen is commonly prepared by the action of an acid, like hydrochloric or sulphuric acid, on certain metals, like zinc or iron, for example. It is a colorless, odorless, tasteless gas, which is very light, in fact the lightest of all known gases, being 14.388 times lighter than air. One liter of pure dry hydrogen at C. and under atmospheric pressure at sea level weighs 0.08987 gram. Hydrogen is an inflammable gas. When pure it may be burned from a jet. The flame is almost colorless, but very hot. Mixtures of air and hydrogen, or oxygen and hydrogen, are explosive, and great care must be taken not to heat such mixtures or to ignite them with a flame or a spark. By cooling hydrogen to 253 C. it may be liquified at atmospheric pressure. Liquid hydrogen is clear and colorless, that is, like water in appearance, but it is only 0.07 as heavy as water. At 259 C. hydrogen solid- ifies, forming white crystals. Hydrogen gas is not poisonous, but animals and plants would die from suffocation when kept in hydrogen, because they require oxygen to breathe. Be- 22 HYDROGEN, OXYGEN, HYDROGEN PEROXIDE 23 cause of its lightness hydrogen is used for filling balloons. Though the gas burns with quite a hot flame, it is not used as a fuel in the pure condition, for this would be too expensive. Hydrogen is, however, one of the constituents of coal gas, natural gas, and water gas, all of which are used as fuels. When hydrogen burns in the air or in oxygen, the product formed is water. When steam is passed over red-hot iron, hydrogen is formed and black oxide of iron, hammer black, is simul- taneously obtained, for at the high temperature the iron FIG. 8. Making hydrogen by passing steam over red-hot iron. unites with the oxygen in the water and thus the hydrogen is set free. Magnesium will similarly liberate hydrogen, and in this case it suffices to immerse that metal in boiling water. Sodium will liberate hydrogen from water rapidly even at room temperature. Potassium acts even more vigorously than sodium. In the case of sodium and po- tassium, the oxides formed pass into solution in the excess of water and form dilute lye. These solutions turn red litmus paper blue, are alkaline to the taste, and feel slippery to the touch. Stronger solutions of lye are caustic and disintegrate the flesh. 24 CHEMISTRY AND DAILY LIFE While hydrogen practically does not occur in nature in the free, i.e. uncombined, condition, it is exceedingly abundant as a constituent of many compounds. As already mentioned, one ninth of the weight of water is hydrogen. All plant and animal tissues contain hydrogen. Crude petroleum and all of its various products contain hydrogen. This element is also found in many rocks. Oxygen occurs in the free state in the air. In fact it forms about 21 per cent of the total volume of the air. Combined with other elements, oxygen is also found in large quantities and very widely distributed over the earth. In fact it is the most abundant of all of the elements, which is readily apparent, for 88.88 per cent of the weight of all water is oxygen, and the latter also forms 44 t 4$ P er cen t of the weight of the solid portions of the earth. In all plants and animals oxygen is combined with hydrogen, carbon, nitrogen, phosphorus, and minor amounts of other elements. Pure oxygen is most readily prepared by heating certain compounds which contain it, and which give up their oxygen content fairly readily at higher temperatures. The substance which is most frequently chosen for this purpose is potassium chlorate. It consists of white crystals whose constituents are potassium, chlorine, and oxygen ; the latter element forms 39.1 per cent of the weight of this salt. Thus from 100 Ib. of potassium chlorate 39.1 Ib. of oxygen may be obtained. The oxides of mercury and silver will also decompose on heating, yielding free oxygen, and metallic mercury and silver respectively. Again, saltpeter when heated will give up a portion of its oxygen. Plants are continually giving off oxygen as they breathe. Pure oxygen is colorless, odorless, and tasteless. It is 1.1 times as heavy as air and sixteen times as heavy as hydrogen. A liter of oxygen at and under a barometric HYDROGEN, OXYGEN, HYDROGEN PEROXIDE 25 pressure of 760 mm. (i.e. under the so-called standard con- ditions of temperature and pressure under which gases are measured) weighs 1.429 grams. Liquid oxygen is pale blue in color, boils at 182. 5 C., and is 1.1315 times as heavy as water. It is attracted by a magnet. By chilling liquid oxygen sufficiently, snow-white crystals may be obtained, which melt at 227 C. In water oxygen is but slightly soluble, 1 volume of water absorbing but 0.034 volume of oxygen at 15 C. This slight solubility of oxygen in water enables us to collect the gas over water. The fact that oxygen is soluble in water, even if but slightly, is nevertheless very im- portant, for fishes and other organisms living in water depend upon this dissolved oxygen for their supply. In its ability to form compounds with other elementary substances oxygen stands at the head of the entire list of the chemical elements. In fact, it will unite with all of the elements except fluorine and the members of the so-called argon group, namely argon, helium, neon, krypton, and xenon. All ordinary combustion in the air is really union of the burning substance with oxygen. That is to say, ordinary combustion is oxidation. So when coal, petroleum, paper, or wood are burned in the air they are oxidized. A part of the products formed is gaseous and so escapes in the air, and another portion remains behind in the solid state as ash. Now, when substances are burned in pure oxygen, the process goes on far more rapidly and more vigorously. In fact, some substances will burn brilliantly in pure oxygen, whereas in the air they either would not burn at all or the action would proceed very slowly, for it must be borne in mind that the oxygen in the air is diluted, so to say, with four times its vol- ume of nitrogen, which is a rather inert gas and does not take part in ordinary combustion. Heated charcoal glows bril- liantly in oxygen. Sulphur and phosphorus burn brilliantly 26 CHEMISTRY AND DAILY LIFE in oxygen gas, and even the steel of a watch spring burns with brilliant scintillations in that gas. One can breathe pure oxygen for a time without bad effects ; in fact, the gas is fre- quently administered to patients who experience difficulty in breathing and so get an insuffi- cient oxygen supply when breathing air. It is in- advisable, however, for a normal person to breathe oxygen, for thus the sys- tem gets too great a sup- ply of that substance and the oxidation processes in the system go on too rapidly. The oxides of carbon, phosphorus, sulphur, and certain other elements readily dissolve in water and form solutions that are sour to the taste and turn blue litmus red. Such solutions are called acid solutions. It was at one time thought that all acids contain oxygen ; this is, however, not the case, though it is true that by far the larger number of all known acids do contain oxygen. The word oxygen means acid generator. It has been retained, even though all acids do not contain oxygen. The fact is that all acids do con- tain hydrogen, which will be considered later. Some oxides, for example, the oxides of sodium, potassium, and calcium (obtained by burning each of these metals in FIG. 9. Burning phosphorus in oxygen. HYDROGEN, OXYGEN, HYDROGEN PEROXIDE 27 oxygen) are white powders which dissolve in water and yield solutions that are alkaline to the taste, slippery to the touch, and turn red litmus blue. Such solutions are called alkaline solutions. When acid solutions and alkaline solutions are mixed, they mutually neutralize each other. The product thus formed is a salt, which, remaining in solution, has neither an acid nor an alkaline taste, and does not affect either red or blue litmus (see acids, alkalies, salts). It thus appears that by union with some elements acidic oxides or acid-form- ing oxides result, and by union with other elements alkaline oxides or alkali-forming oxides are obtained. It is true also that in some cases oxides are formed which are neither acid nor alkali forming, but rather indifferent or neutral bodies, like the oxides of iron, copper, and mercury, for example. Ozone may be prepared from oxygen. Three volumes of oxygen yield two volumes of ozone. On heating the latter ordinary oxygen is formed again, three volumes being obtained from two volumes of ozone. Ozone is then 1.5 times as heavy, as oxygen. Ozone is produced when pure oxygen or the air is subjected to the action of electric sparks. So, for instance, in the neighborhood of a frictional electrical machine in action there is always a certain amount of ozone produced, readily distinguished by its peculiar garlic-like odor. This odor of ozone is also observed when lightning strikes terrestrial objects or when phosphorus slowly oxidizes at room temperatures in moist air. Ozone is the most powerful oxidizing agent known and this is its most important property. Ozone will not only kill germs and other microscopic organ- isms rapidly, but it will also oxidize dyestuffs, destroying their color, and convert silver, lead, arsenic, sulphur, and other elements into oxides. On account of its great activity, ozone exists in the air for but a short time after it has been formed. Ozone acts on water, forming hydrogen peroxide, 28 CHEMISTRY AND DAILY LIFE also called hydrogen dioxide. This is a compound of oxygen and hydrogen which contains just twice as much oxygen as does water. Thus in water 1 gram hydrogen is united to 8 grams of oxygen, while in hydrogen peroxide 1 gram of hydrogen is united to 16 grams of oxygen. In the pure state hydrogen peroxide is a thick sirupy liquid whose specific gravity is 1.458. It is colorless, but like water it looks blue in thick layers. This liquid is quite unstable and decomposes readily into water and oxygen. In dilute solutions it is more stable, especially when kept cool and in the dark. It is now a common article of commerce, the 3 per cent solutions frequently being sold under the trade name of " dioxygen." This solution is valuable as a mild bleaching agent and as an antiseptic. Its power as a germicide depends upon the fact that it readily decomposes into oxygen and water. The oxygen set free destroys the organisms. Hydrogen peroxide is commonly prepared by the action of cold dilute sulphuric acid upon barium peroxide, thus : Barium peroxide plus sulphuric acid yields hydrogen peroxide plus barium sulphate. The latter substance may be filtered off. The clear filtrate contains the hydrogen peroxide in solution. On account of the fact that hydrogen peroxide also results when water is acted upon by ozone, the latter cannot exist long in the air, which always contains moisture, and this would interact with ozone, forming hydrogen peroxide and oxygen. QUESTIONS 1. What are the important properties of hydrogen ? 2. How may hydrogen be prepared ? 3. In what compounds is hydrogen found in nature ? 4. What is formed when hydrogen burns? How may this be shown ? HYDROGEN, OXYGEN, HYDROGEN PEROXIDE 29 5. How does oxygen occur in nature ? 6. How may pure oxygen be made ? 7. How many pounds of oxygen could be prepared from 83 pounds of potassium chlorate ? 8. Mention the important properties of oxygen. 9. What is an oxide ? How may oxides of phosphorus, carbon, sulphur, and iron be formed ? 10. How may ozone be formed ? Mention its important proper- ties. What is ozone ? 11. What is hydrogen peroxide ? How may it be prepared ? 12. Mention the uses of hydrogen peroxide. CHAPTER IV THE AIR, NITROGEN, NITRIC ACID, AND AMMONIA THE main constituents of the air are nitrogen and oxygen. These are present in about the proportions of 78 volumes of the former to 21 volumes of the latter. Besides these two gases there are present, however, water vapor, carbon dioxide, and the elements of the argon group, namely helium, neon, krypton, and xenon. Table 5 gives the composition of a sample of air as it is found in the country or over the sea. TABLE 5 COMPOSITION OF A SAMPLE OF NORMAL COUNTRY AIR 100 volumes contain as follows : Nitrogen 77.42 volumes Oxygen 20.77 volumes Argon, helium, neon, krypton, and xenon . . 0.93 volume Water vapor 0.85 volume Carbon dioxide 0.03 volume 100.00 volumes The amount of moisture which air contains varies consider- ably from time to time, depending especially upon temperature and proximity to bodies of water. Minor quantities of sulphur dioxide, hydrogen, ammonia, nitric acid, particles of dust, and various microbes are also present in air. The amounts of all of these are also quite variable. While the members of the argon group are present in air to the extent of over 0.9 per cent by volume, these gases are nevertheless of no practical importance and hence will receive no further con- 30 AIR, NITROGEN, NITRIC ACID, AMMONIA 31 sideration here. The carbon dioxide in the air in the country or over the sea amounts to about 0.03 per cent by volume, but in cities where much fuel is consumed the air often contains double that amount and even more, while in densely crowded audience rooms the carbon dioxide content may run as high as six to eight times that found in city air. While the carbon dioxide in the air is present to the extent of only about 0.03 per cent, yet it must be borne in mind that plants get all of their carbon from the carbon dioxide in the atmosphere. They breathe in carbon dioxide, which in the presence of sunlight in the green leaf of the plant interacts with the moisture that is present, forming starch and setting oxygen free, which the plant therefore exhales. This change may be written as follows : Carbon dioxide + water (in the sunlight in the green leaf of the plant) yields starch + oxygen. The constituents of the air, Table 5, are not chemically united; they form merely a mixture. In spite of this the relative amounts of oxygen, nitrogen, and argon are nearly constant in the atmosphere everywhere. The highest amounts of moisture which the air is able to hold at different temperatures is shown in Table 6. TABLE 6 TEMPERATURE OF AMOUNT OF WATER VAPOR IN SATURATION 1 Cu. METER OF SATURATED AIR -5 C 4.0 grams 0C 5.4 grams +5 C ; . . 7.3 grams 10 C 9.7 grams 15 C 13.0 grams 20 C 17.1 grams 22 C 22.5 grams 30 C. 30.0 grams 32 CHEMISTRY AND DAILY LIFE If the air has but 0.4 of the water vapor in it that it can hold, we feel that the air is dry. If, however, 0.8 of its maximum capacity of moisture is present in the air, we feel that it is moist or humid. As the moist air reaches the upper and cooler regions of the atmosphere, partial con- densation of the water to minute drops takes place, form- ing mists, which on account of their altitude are commonly called clouds. When the mists form near the surface of the earth, they are called fogs. These are formed when mois- ture-laden air is cooled to such an extent that a portion of its moisture is condensed. This is frequently caused by a cold current of air striking warmer moisture-laden strata. When the condensation in the upper regions of the air is such that large drops are formed, these fall to the ground as rain, and carry with them in solution some of all of the gases in the atmosphere, together with fine particles of dust, microor- ganisms, etc., that were suspended in the air. This material on reaching the ground soaks into the soil in part, and in part it flows off the surface into brooks, rivers, lakes, and the sea. It has been estimated that about 80 per cent of the water that falls as rain on the land soaks into the ground and 20 per cent runs off the surface into the waterways. Snow, sleet, and hail form at temperatures at and below the freezing point of water. Dew forms at night. It consists of drops of moisture that deposit upon objects which cool off relatively rapidly, and so chill the moisture-laden air that comes into contact with them. The importance of these various aqueous precipitations for the life of plants will receive further consideration later. The common way of preparing nitrogen gas is to abstract the oxygen from an inclosed space of air. This may be done in various ways. So, for example, an animal like a rat or a AIR, NITROGEN, NITRIC. ACID, AMMONIA 33 mouse may be placed in a tightly stoppered bottle, where- upon it will soon suffocate, because it will use up the oxygen in the air in the bottle and breathe out carbon dioxide. The latter gas may be removed by shaking the gas left in the bottle with some clear limewater. The limewater be- comes milky in appearance because a precipitate of car- bonate of lime, calcium carbonate, is formed, thus : Limewater -f- carbon dioxide yields calcium carbonate + water. By this chemical change, then, the carbon dioxide is removed, being made part and parcel of a solid substance, the calcium carbonate, which is chemically the same as chalk. Now the gas which still remains in the bottle is nitrogen plus the gases of the argon group. These latter in all make up only about 0.9 per cent of the entire air. They have all been discovered in the air since 1895. They are very inert, in fact thus far it has been found to be impossible to get them to unite chemically with any other element. By simply removing the oxygen and carbon dioxide from the air, pure nitrogen is not obtained, for the gases of the argon group are still present. For the present purposes, however, it will not be necessary to devote space to the description of how the latter gases may be removed; suffice it here to say that perfectly pure nitrogen is best prepared from certain pure chemical compounds in which it occurs. The oxygen of the air may also be removed by burning certain substances in the air in a closed space. So one may burn a bit of sponge FIG. 10. Removing oxygen from the air by burning alcohol in it. 34 CHEMISTRY AND DAILY LIFE saturated with alcohol in a bottle inverted in a dish of water as shown in Fig. 10. In this case water and carbon dioxide are formed and the latter gas is absorbed by the water in the bottle. In place of the alcohol a bit of phos- phorus may be burned similarly in the bottle. In this case the phosphorus unites with the oxygen, forming phosphoric oxide, which dissolves in the water, thus leaving the nitro- gen intact. In both of the last experiments the water rises on the inside of the bottle, filling the latter to the extent of one fifth of the original volume occupied by the air, thus showing that four fifths of the air is nitrogen and one fifth oxygen. If a burning splinter or a lighted candle is thrust into nitrogen gas, the flame is at once extinguished, showing that nitrogen will not support combustion. Animals suffocate in nitrogen, and die for lack of oxygen. Nitrogen gas is not poisonous, but it cannot be used in respiration. In the air it dilutes the oxygen that is present and so retards the too rapid oxidation that would take place if pure oxygen were inhaled. At ordinary temperatures nitrogen is quite an inert gas, but at high temperatures it does combine with other sub- stances chemically. So, for example, nitrogen does not unite with oxygen when these gases are mixed. In fact no union takes place, even if these gases are heated together quite highly, as, for instance, when the mixture is passed through a red-hot tube. But if a mixture of nitrogen and oxygen is subjected to the very high temperature of the electric arc, union does take place, oxide of nitrogen being thus formed. The latter when treated with air and water yields nitric acid, and indeed this process is now used successfully in form- ing that important substance on a commercial scale. As electricity is required in the process, this method of manu- facturing nitric acid is only profitable where water power is AIR, NITROGEN, NITRIC ACID, AMMONIA 35 available for the production of cheap electric energy. In Norway considerable quantities of nitric acid and nitrates, which are important as fertilizers, are thus manufactured at present. As the lightning flashes through air during thunder- storms, nitric acid and nitrates are similarly produced, and these as they are carried down by the rain enrich the soil. It is for this reason that a thundershower is much more helpful to vegetation than merely an ordinary rain. The latter supplies water, but the former furnishes fertilizer as well as water. The two most important common compounds of nitrogen are nitric acid and ammonia. How nitric acid may be made FIG. 11. Making nitric acid. from the air and water has just been stated. It may also be made from saltpeter, i.e. potassium nitrate, by treatment with strong sulphuric acid and heating. Thus the nitric acid is liberated and distills over into the receiver as shown in Fig. 11. Instead of the somewhat costly potassium ni- 36 CHEMISTRY AND DAILY LIFE trate, sodium nitrate, which is mined in Chili and conse- quently called Chili saltpeter, may be used. Thus nitric acid and sodium sulphate are produced, instead of nitric acid and potassium sulphate. We may write these changes thus : Potassium nitrate 4- sulphuric acid = nitric acid 4- potassium sulphate. Sodium nitrate + sulphuric acid = nitric acid + sodium sulphate. Pure nitric acid is a colorless liquid having a pungent odor. Its specific gravity is 1.414 at 15 C. In the sunlight the liquid turns yellowish in color, due to partial decomposition. Nitric acid is a wry powerful acid and also a strong oxidizing agent. It turns blue litmus red, discolors indigo solution, disintegrates cloth, corrodes the flesh, and turns the skin yellow. Metals, with the exception of gold and platinum, are acted upon by nitric acid, being converted either into nitrates or oxides. The various salts of nitric acid are all called nitrates. The nitrates, especially those of potassium, ammonium, calcium, and sodium, are very important as fertilizers. While nitric acid does not occur in nature as such, its salts, that is, the nitrates, are quite widely distributed in the soil. In the air nitrate of ammonium is also present. Nitrates are exceedingly important as soil constituents, for it is from them that plants get their supply of nitrogen. The nitrates in the soil are contained in the soil water which holds them in solution. Nitric acid is also used in manufacturing sulphuric acid, in which case it serves to oxidize sulphur to its highest stage of oxidation. Furthermore, in making collodion, gun cotton, dynamite, and smokeless powder, nitric acid is used in large quantities. AIR, NITROGEN, NITRIC ACID, AMMONIA 37 Nitrous acid may be obtained from nitric acid by robbing the latter of a part of its oxygen. The salts of nitrous acid are called nitrites ; they occur in the soil, as a result of the decomposition of plant and animal matter. All plants and animals contain nitrogen in combination with carbon, hydrogen, oxygen, and minor amounts of sulphur and phos- phorus. When the plant and animal tissues decay, nitrogen is set free in part, and in part it is converted into ammonia, a compound of nitrogen and hydrogen, and the ammonia is gradually oxidized to nitrites and finally to nitrates. The latter represent the highest oxidation products. While the nitrates in the soil are a very important source from which plants derive their supply of nitrogen, they are by no means the only source. Ammonia is a compound of nitrogen and hydrogen. It contains 14 parts of the former to 3 parts of the latter by weight. By volume the composition is represented by the following equation : 3 volumes of hydrogen 4- 1 volume of nitrogen = 2 volumes of ammonia. Hydrogen and nitrogen do not combine at all readily, how- ever, even when the gases are heated together highly, as, for instance, by means of the electric spark. On thus subject- ing a mixture of hydrogen and nitrogen gases to the contin- uous action of the electric spark, a small fraction of a per cent of the gases present is converted to ammonia. But by heating nitrogenous plant or animal tissues, with little or no access of air, ammonia is readily obtained, especially when the organic matter used is mixed with lime or caustic soda or potash and then subjected to heat. In fact the ammonia and all of the ammonium compounds of commerce are obtained as a by-product of the manufacture of illuminating gas from 38 CHEMISTRY AND DAILY LIFE coal. In the latter process the coal, which represents the remains of the vegetation of the carboniferous age and other early geological periods, is heated out of contact of the air. The gases that escape are washed by gurgling them through water, and the ammonia is then found dissolved in this water. Plant and animal matter, including coal, peat, etc., all contain nitrogen and hydrogen in combination with carbon, oxygen, sulphur, phosphorus, etc., and when this material is heated out of contact with the air (i.e. subjected to dry distillation, also called destructive distillation), a variety of gaseous products are formed, among which is ammonia. The latter forms by union of nitrogen and hydro- gen from the organic matter. Ammonia gas is colorless and only 0.59 as heavy as air. It has a very characteristic, sharp, penetrating odor, irri- tates the mucous membranes, and causes a flow of tears. One cannot inhale the gas. Animals soon die in it. The gas does not support combustion and does not burn in the air ; but it may be burned in oxygen. About 700 volumes of ammonia gas will dissolve in 1 volume of water at ordi- nary temperatures. This solution is lighter than water. At 14 C. the saturated solution has a specific gravity of 0.8844 and contains 36 per cent of ammonia by weight. The solution of ammonia gas in water is popularly called ammonia water, spirits of hartshorn, and aqua ammonia. It has the same odor as ammonia gas, for the latter is con- tinually being given off by the solution. Indeed, all of the ammonia gas may be expelled from the solution by boiling the same. Finally water alone remains behind. Ammonia turns red litmus blue, and is consequently alkaline, which is further evidenced by the fact that it will unite directly with acids, neutralizing them and forming ammonium salts. So with nitric acid ammonia forms ammonium nitrate, with AIR, NITROGEN, NITRIC ACID, AMMONIA 39 sulphuric acid ammonium sulphate, with hydrochloric acid ammonium chloride. All ammonium salts may be volatilized by heating them. On treating any ammonium salt with lime or lye, ammonia gas is evolved, and this is the best way to prepare pure ammonia gas. For example : Ammonium chloride H- soda lye = sodium chloride + ammonia + water. (common salt) Ammonia water is used in the household in cleansing clothes and polishing metals. It does not attack the latter as drastically as acids do, and consequently is less objectionable. Ammonia is formed wherever nitrogenous plant and animal remains are decaying. So, for instance, it is found in stables, in dung piles, in the leachings from the latter, and in the soil. When thus formed from rotting organic matter, ammonia at once combines with any acids that may be present, especially with carbonic acid gas, which is always at hand, being a constituent of the air. In the soil, then, ammonia is present in the form of ammonium salts, among these ammonium nitrate, ammonium nitrite, ammonium car- bonate, ammonium chloride, and ammonium sulphate are the most important. All of the ammonium salts are soluble in water, and hence are present in solution in the soil water. From ammonium salts thus dissolved in the waters of the soil, plants derive a considerable share of their supply of nitrogen, hence the value of ammonium salts as fertilizers. In com- merce ammonium sulphate, which contains 21 per cent nitro- gen, is quite generally sold as a fertilizer. It is manufac- tured at the gas works in large cities as a by-product in making coal gas. 40 CHEMISTRY AND DAILY LIFE QUESTIONS 1. (a) What are the constituents of the air ? (b) In what proportions do the two main constituents occur in the air ? 2. Of what use is the carbon dioxide in the air ? 3. When does the air feel dry ? When moist ? 4. What becomes of the water that falls to the ground as rain ? 6. What is dew ? Snow ? Sleet ? Hail ? 6. How may nitrogen be prepared ? Give its principal proper- ties. 7. Why will an animal die in nitrogen gas ? 8. How may nitric acid be formed ? 9. What are the properties of nitric acid ? 10. Of what use are nitrates in the soil ? 11. What is ammonia ? How prepare it ? 12. What is ammonia water ? What is it used for ? 13. Why are ammonium salts useful as fertilizers ? From what source are these salts obtained ? 14. Mention the important forms in which nitrogen occurs in nature. CHAPTER V ACIDS, ALKALIES, SALTS, AND CHEMICAL FORMULAS WE have already learned that when carbon, sulphur, or phosphorus are burned and the products of combustion are dissolved in water, liquids are obtained which have a sour taste and redden blue litmus. Substances possessing such characteristics are commonly termed acids. Acids occur in nature in plants, in animals, and also in the mineral world. So, for example, the sourness of the lemon is due to the citric acid which it contains. The latter substance may be prepared from lemon juice. It consists of beautiful color- less crystals which are readily soluble in water. The solu- tion has a markedly sour taste. This same citric acid is found in other citrous fruits as well, and is most abundant in them before they are quite ripe. In sour apples, moun- tain ash berries, and many other similar fruits malic acid is present. It too may be prepared from these in the pure state. It is a white solid substance which is readily soluble in water. Similarly grapes contain tartaric acid. The jack- in-the-pulpit contains oxalic acid. On fermentation of apple juice, vinegar, which is essentially a dilute solution of acetic acid, is formed. When milk sours, lactic acid is formed, and indeed lactic acid also occurs in the muscles of man and animals. In the human stomach hydrochloric acid is formed by the gastric glands. In red ants formic acid is found. The air, as we have seen, always contains a small amount of carbonic acid, and in cities where much coal is consumed, it 41 42 CHEMISTRY AND DAILY LIFE also contains sulphurous acid, for coal contains sulphur.. All natural waters, too, contain carbonic acid in solution, and some spring waters are highly charged with this sub- stance. Soil waters also contain acids formed during the processes of decay of animal and vegetable matter. Boric acid is found in volcanic regions to some extent, and silicic acid is one of the most abundant of all substances. It occurs as quartz in huge masses, often forming mountains, and as sand grains it covers vast areas of the earth's surface. Silicic acid is practically insoluble in water. It conse- quently has no taste and does not redden litmus. How- ever, with alkalies it has the power to form salts, which indeed is after all the most important characteristic of all acids. In fact some acids are so weak that they do not redden litmus, and haw no perceptible sour taste, and yet they are unquestionably acidic substances because they do form salts with alkalies. Lime, as it is used for building purposes, soda lye, potash lye, and ammonia are substances which turn moist red litmus paper blue. They are typical alkalies. In strong solutions they have a caustic or corrosive action on .the skin. With acids they react chemically, forming new substances called salts. So, for example, when a solution of soda lye is treated with a solution of hydrochloric acid till the resulting liquid does not change either red or blue litmus paper, we say that the acid and the lye have neutralized each other. The sub- stance which is now in the solution is common salt. It has a salty taste, whereas the acid had a sour taste, and the lye had an alkaline taste which is quite peculiar to itself. The change which takes place when the acid and the lye act on each other may be expressed in the form of an equation thus : Soda lye + hydrochloric acid yields sodium chloride (i.e. common salt) + water. ACIDS, ALKALIES, SALTS, CHEMICAL FORMULAS 43 Now it is somewhat cumbersome to write such equations in words, and in order to save space and time, chemists have adopted a system of abbreviations which is very easily comprehended. Each element is represented by the initial letter of its name. Thus C is the symbol for carbon, H stands for hydrogen, O for oxygen, N for nitrogen, P for phos- phorus, I for iodine, etc. As several of the elements have names that begin with the same letter, their symbols are distinguished from one another by adding to the first letter (which is always capitalized) one other characteristic letter contained in the name of the element. This second letter is never capitalized. So, for example, C stands for carbon, Ca for calcium, Co for cobalt, Cl for chlorine, etc. While in some cases the symbol is derived from the common name of the element (as already illustrated), in other instances it is derived from the Latin name of the element. So sodium, natrium, has the symbol Na ; potassium, kalium, K ; copper, cuprum, Cu ; silver, argentum, Ag ; iron, ferrum, Fe ; mer- cury, hydrargyrum, Hg ; tin, stannum, Sn ; gold, aurum, Au ; (compare Table 1). Now these symbols do not stand for the names of the elements only, but they also represent definite quantities by weight in which the elements combine chemically. So H stands for hydrogen, but also for 1 gram of hydrogen ; Cl stands for chlorine, but also for 35.5 grams of chlorine ; O represents oxygen, but also 16 grams of oxy- gen ; etc. Table 7 gives the names of the more common elements together with their symbols and the number of parts by weight for which each symbol stands. A complete list of the chemical elements and their symbols has already been given on page 3. While Table 7 mentions but thirty- eight of the elements, the data are nevertheless quite suffi- cient for all ordinary purposes, for the elements that have been omitted are of minor importance. 44 CHEMISTRY AND DAILY LIFE Al 27.1 Lead Pb 207.10 Sb 120.2 Lithium . . . . Li 6.94 As 74.96 Magnesium . . . Mg 24.32 Ba 137.37 Manganese . . . Mn 54.93 Bi 208.0 Mercury . . : Hg 200.6 B 11.0 Molybdenum . . Mo 96.0 Br 79.92 Nickel . . . Ni 58.68 Cd 112.40 Nitrogen . . . N 14.01 Ca 40.07 Oxygen . . . . 16.0 C 12.0 Phosphorus . P 31.04 Cl 35.46 Platinum . . . Pt 195.2 Cr 52.0 Potassium . . . K 39.10 Co 58.97 Silicon . . . . Si 28.3 Cu 63.57 Silver . . . . Ag 107.88 F 19.0 Sodium . . . . Na 23.0 Au 197.2 Strontium . . . Sr 87.62 H 1.008 Sulphur . . . . S 32.07 I 126.92 Tin .... . Sn 119.0 Fe- 55.84 Zinc . Zn 65.35 TABLE 7 Aluminum . . Antimony Arsenic . . Barium . . Bismuth . . Boron . . . Bromine . . Cadmium . . Calcium . . Carbon . . Chlorine . . Chromium Cobalt . . . Copper . . . Fluorine . . Gold . . . Hydrogen . . Iodine . . . Iron . . . The figures in Table 7 have all been determined by ascertaining the weights of the elements that unite to form compounds. Experience has shown that chemical compounds when pure always contain the same elements in the same pro- portion by weight. This is called the law of definite propor- tions. So, for instance, hydrochloric acid always consists of hydrogen and chlorine, and nothing else. Moreover, in hydrochloric acid there are 35.5 grams of chlorine combined with every gram of hydrogen; that is to say, for every 1 part by weight of hydrogen, hydrochloric acid contains 35.5 parts by weight of chlorine, and this is represented by the symbol HC1, commonly termed the formula for hydro- chloric acid. Again, in common salt, which is sodium chlo- ride, we always have 35.5 parts by weight of chlorine com- bined with 23 parts by weight of sodium. The symbol or formula for common salt, .then, is NaCl, and it expresses both the qualitative and quantitative composition of sodium chloride. ACIDS, ALKALIES, SALTS, CHEMICAL FORMULAS 45 We may indeed regard common salt, i.e. NaCl, as derived from hydrochloric acid, i.e. HC1, the 1 part by weight of hydrogen of the latter being replaced by 23 parts of sodium by weight. Thus 23 parts of sodium by weight may replace 1 part of hydrogen by weight, and consequently 23 is said to be the hydrogen equivalent of sodium ; for it has been found to be true in general that whenever sodium can replace hydrogen in a compound it takes 23 grams of sodium to play the role of 1 gram of hydrogen. The hydrogen equivalent of other elements may be found similarly. Such replacements are very easily represented by means of the chemical symbols mentioned. Thus : HC1 + Na (36.5 grams hydrochloric acid) + (23 grams sodium) yield NaCl + H (58.5 grams common salt) + (1 gram hydrogen) In general, then, the composition of any compound is expressed by writing the symbols of its elements one after the other. How- ever, whenever the compound is a gas or vapor, the symbols that represent it indicate^ not only the qualitative composi- tion and quantitative composition by weight, but also the weight of 22.4 liters of the gas or vapor at C. and 760 mm. barometric pressure. So, for example, the symbol HC1 stands for hydrochloric acid and indicates that in this com- pound 1 gram of hydrogen is combined with 35.5 grams chlorine ; it shows also, however, that 22.4 liters of the gas or r weigh 36.5 grams, whence 1 liter weighs ^rr> or 1-629 grams. 22.4: The volume 22.4 liters is chosen because it is the volume of 2 grams of hydrogen at C. and 760 mm. pressure, with which it is customary to compare other gases. It is not necessary to enter further into the reasons for this practice here. For our purpose it suffices to know that in the case 46 CHEMISTRY AND DAILY LIFE of any gas the weight of a liter of it may be found by dividing the formula weight by 224- The following examples will serve as illustrations. The composition of carbon dioxide gas is represented by CO 2 , which indicates that 12 grams of carbon are combined with 2X16 or 32 grams of oxygen ; furthermore, the formula indicates that 12 + 32 or 44 grams (the formula weight) is the weight of 22.4 liters of the gas, 44 whence 1 liter weighs rr-r, or 1.964 grams. Again, the for- A mula for water is H 2 O, which shows that it consists of hydro- gen and oxygen in the ratio of 2 grams of the former to 16 grams of the latter; it also indicates, however, that 2 + 16 or 18 grams of water vapor occupy 22.4 liters of space under standard conditions, whence 1 liter of the vapor I O weighs on~i> or 0.8035 gram. From the above illustrations the use of subscripts to the symbols of the elements is per- haps already sufficiently evident, but two further illustra- tions will here be given. The symbol for cane sugar is C^H^Oii, which shows that this compound is composed of 12 X 12 or 144 parts of carbon by weight, to every 1 X 22 or 22 parts of hydrogen by weight, to every 11 X 16 or 176 parts of oxygen by weight ; in other words every 342 grams of sugar contain 144 grams of carbon plus 22 grams of hydro- gen plus 176 grams of oxygen. The symbol for phosphoric acid anhydride, also called phosphorus pentoxide, is P2O 5 . This formula shows that this compound is composed of phos- phorus and oxygen in the proportions of 2 X 31 or 62 grams of the former to 5 X 16 or 80 grams of the latter. The use of chemical symbols, then, serves to express quite a number of facts about a chemical compound in compact form. A few of 'the chemical changes already studied will now be expressed, using chemical symbols and equations. ACIDS, ALKALIES, SALTS, CHEMICAL FORMULAS 47 These will serve to illustrate further how the formulas are used. (1) When iron filings and sulphur are heated together, ferrous sulphide is formed. Thus : Fe + S = FeS iron sulphur ferrous sulphide 56 grams 32 grams 88 grams (2) When carbon is burned in oxygen or in the air, car- bon dioxide is formed. Thus : C + 2 = CO 2 carbon oxygen carbon dioxide 12 grams 32 grams 44 grams (3) When phosphorus is burned in oxygen, phosphorus pentoxide is formed. Thus : P 2 + 5 = P 2 6 phosphorus oxygen phosphorus pentoxide 62 grams 80 grams 142 grams (4) When sulphuric acid acts on zinc, hydrogen and zinc sulphate are formed. Thus : H 2 SO 4 + Zn = H 2 + ZnS0 4 sulphuric acid zinc hydrogen zinc sulphate 98 grams 65 grams 2 grams 161 grams (5) When hydrochloric acid is neutralized by sodium hydroxide, sodium chloride and water are formed. Thus : HC1 + NaOH = NaCl + H 2 O hydrochloric acid sodium hydroxide sodium chloride water 36.5 grams 40 grams 58.5 grams 18 grams (6) When potassium chlorate is heated, it yields oxygen and potassium chloride. Thus : KC1O 3 3O + KC1 potassium chlorate oxygen potassium chloride 122.5 grams 48 grams 74.5 grams 48 CHEMISTRY AND DAILY LIFE There are really only three different types of chemical change possible, namely : (1) When two or more substances unite to form one new substance, as in reactions (1), (2), and (3) above. These are all examples of synthesis. (2) When two or more substances are formed by decom- position of a single substance, as in reaction (6) above. This may be termed analysis, as contrasted with synthesis. (3) When two or more substances react with one another to form two or more new substances, as in reactions (4) and (5). These are also termed cases of double decomposition. Chemical symbols, then, are a great aid in that they indi- cate the chemical composition of a compound at a glance. It must, of course, be kept in mind that such symbols are only the expression of facts found out by experiment, and so the formula of a compound cannot be written till after its com- position has been actually ascertained by chemical analysis. The formation of salts really appears much simpler when expressed by means of equations in which chemical symbols are used, thus : KOH + HNO 3 = KNO 3 + H 2 O potassium hydroxide nitric acid potassium nitrate water or potash lye or saltpeter 2 NaOH + H 2 SO 4 = Na 2 SO 4 + 2 H 2 O sodium hydroxide sulphuric acid sodium sulphate water or soda lye Ca(OH) 2 + 2 HC1 = CaCl 2 + 2 H 2 O calcium hydroxide hydrochloric calcium chloride water or slaked lime acid In regarding the last three equations it will be seen that an acid is a compound containing hydrogen which may be replaced by a metal, thus forming a salt. So potassium ni- trate, sodium sulphate, and calcium chloride are typical ACIDS, ALKALIES, SALTS, CHEMICAL FORMULAS 49 salts. Now the hydroxides of potassium, sodium, and calcium are typical bases. A base is a compound (usually an hydroxide of a metal) which upon reacting with an acid forms a salt and water. And finally a salt is a compound formed when a base acts upon an acid. We may consider salts as the products formed when the hydrogen of an acid is replaced by a metal or some group of elements that may play the role of a metal so far as the process of a salt formation is concerned. Common salt, NaCl, may be formed by direct union of sodium and chlorine, thus : Na + Cl = NaCl or by neutralization of the base, sodium hydroxide, NaOH, by hydrochloric acid, thus : NaOH + HC1 = NaCl + H 2 O Calcium sulphate, CaS0 4 , may be formed by the union of lime, CaO, with sulphuric anhydride, SO 3 , thus : CaO + SO 3 = CaSO 4 or by the action of slaked lime on sulphuric acid, thus : Ca(OH) 2 + H 2 S0 4 = CaS0 4 + 2 H 2 O The metals, then, are in general base-forming substances, and the non-metals are acid-forming substances, though some metals may at times act as acid-forming elements, and some of the non-metals may exhibit basic properties. Among the most common salts are chlorides, sulphates, nitrates, carbonates, phosphates, and silicates. These may be considered as derived respectively from the following acids: hydrochloric acid HC1, sulphuric acid H 2 SO 4 , nitric acid HNO 3 j carbonic acid H 2 CO 3 , phosphoric acid H 3 PO 4 , 50 CHEMISTRY AND DAILY LIFE and silicic acid H 2 SiO 3 . So, for example, we have sodium chloride NaCl, calcium chloride CaCl 2 , potassium chloride KC1, magnesium chloride MgCl 2 , ferric chloride FeCl 3 , cupric chloride CuCl 2 , etc., all of which may be considered as derived from hydrochloric acid HC1, the hydrogen of which is replaced by the respective metals. Again, we have sodium sulphate Na 2 SO 4 , potassium sulphate K 2 SO 4 , am- monium sulphate (NH 4 ) 2 SO 4 , copper sulphate CuSO 4 , cal- cium sulphate CaSO 4 , magnesium sulphate MgSO 4 , ferrous sulphate (also commonly called copperas) FeSO 4 , etc. These sulphates may all be considered as derived from sulphuric acid H 2 SO 4 , whose hydrogen has been replaced by the metals or the ammonium group respectively. Similarly we have potassium nitrate KNO 3 , sodium nitrate NaNO 3 , calcium nitrate Ca(NO 3 ) 2 , copper nitrate Cu(NO 3 ) 2 , ammonium nitrate NH 4 NO 3 , etc., all of which may be regarded as derived from nitric acid HNO 3 . The carbonates like so- dium carbonate Na 2 CO 3 , potassium carbonate K 2 CO 3 , ammonium carbonate (NH) 2 CO 3 , calcium carbonate CaCO 3 , lead carbonate PbCO 3 , etc., may be considered as derivatives of carbonic acid H 2 CO 3 . The phosphates like calcium phosphate Ca 3 (PO 4 ) 2 , sodium phosphate Na 2 HPO 4 , potas- sium phosphate K 2 HPO 4 , ferric phosphate FePO 4 , may all be regarded as derived from phosphoric acid H 3 PO 4 ; whereas silicates like sodium silicate Na 2 SiO 3 , calcium silicate CaSiO 3 , etc., may be considered as derivatives of silicic acid. In like manner acetates are derived from acetic acid, arsenates from arsenic acid, lactates from lactic acid, tartrates from tartaric acid, oleates from oleic acid, borates from boric acid, and so on. Each acid is, then, capable of forming a series of salts which result when the hydrogen of the acid is replaced by the various metals or radicals (i.e. groups of ele- ments like ammonium, NH 4 ) that may act as metals. In turn ACIDS, ALKALIES, SALTS, CHEMICAL FORMULAS 51 each metal may form a series of salts as that metal replaces the hydrogen of the various acids. So, for instance, we have the chloride of copper, the nitrate of copper, the sulphate of copper, the borate of copper, the oleate of copper, the phos- phate of copper, the silicate of copper, etc. Each metal may in general form a similar long list of salts with the various acids. QUESTIONS 1. Define the terms acid, base, salt, and give an example of each. 2. Mention six acids found in nature, stating where they occur. 3. What is an alkali ? Give four examples. 4. Write the equation expressing the neutralization of hydro- chloric acid by caustic soda. 5. Explain the meaning of the following symbols : H, 0, Co, P, Cl, N, K, Hg, Au, Pt, S. 6. What is the law of definite proportions ? 7. What is the meaning of the following formulas : H 2 0, NaCl, P 2 5 , C 12 H 22 O n ? 8. How much ferrous sulphide, FeS, may be formed from 200 grams of sulphur ? How many pounds would this be ? 9. How much common salt, NaCl, would be obtained when 2 pounds of sodium are burned in chlorine ? 10. How much lime, CaO, would be required to produce 1000 grams of chalk, CaCOs ? CHAPTER VI THE HALOGENS t THE elements fluorine, chlorine, bromine, and iodine are called the halogens. They never occur in nature in the uncombined state, but they are frequently met in compounds. Fluorine is found chiefly as calcium fluoride, fluor spar, CaF 2 , but it is after all fairly widely distributed in minute amounts in granitic rocks, being a minor constituent of the mineral apatite, which consists essentially of calcium phos- phate-plus a small percentage of calcium fluoride. Fluorine is present in extremely small quantities in soils, from which it gets into plants, and so into animals. In the enamel of the teeth of the latter notable amounts of fluorine are always present. By treating calcium fluoride with sulphuric acid hydrogen fluoride, hydrofluoric acid, HF, is formed, thus : CaF 2 + H 2 SO 4 = CaSO 4 + 2 HF calcium sulphuric calcium hydrofluoric fluoride acid sulphate acid This acid is a gas which readily dissolves in water. It attacks glass, sand, and silicates in general, forming silicon fluoride, which is volatile, and metallic fluorides with any bases that may be present. Thus its action on calcium silicate is shown by the following equation : CaSiO 3 + 6HF = CaF 2 + SiF 4 + 3 H 2 O calcium hydrofluoric calcium silicon water silicate acid fluoride tetrafluoride 52 THE HALOGENS 53 Hydrofluoric acid is poisonous. The gas and also the aque- ous solution of the latter are used for etching glass, dia- mond ink. Hydrofluoric acid is also used in the laboratory for the purpose of decomposing silicates in chemical analysis. Fluorine itself may be obtained by passing the electric current through a solution of potassium fluoride and hydro- fluoric acid in water. Fluorine is a light-colored greenish yellow gas of extremely pungent odor. It unites with all elements directly, except with oxygen. It is probably the most active of all of the chemical elements. Chlorine is found mainly in common salt, sodium chloride, XaCl, the occurrence of which will be discussed in connection with sodium. By treating common salt with sulphuric acid, hydrochloric acid is formed, thus : NaCl + H 2 SO 4 = NaHSO 4 + HC1 sodium sulphuric sodium acid hydrochloric chloride acid sulphate acid Hydrochloric acid, also called muriatic acid or spirit of salt, is a colorless pungent gas, which fumes in the air, for it has a strong affinity for water and condenses the latter from the air to drops. Thus the fumes are really minute droplets of a solution of hydrochloric acid gas in water. At ordinary temperatures and atmospheric pressure, one volume of water will dissolve about four hundred volumes of hydro- chloric acid gas. Hydrochloric acid is a powerful acid. With bases it reacts, forming chlorides and water. In the human stomach there is found a 0.33 per cent solution of hydrochloric acid, which aids in the digestion of food. Hydrochloric acid is made in large quantities on a com- mercial scale in connection with the LeBlanc soda process (which see). The saturated aqueous solution has the specific gravity 1.19 and contains 38 per cent of pure hydrochloric acid by weight. 54 CHEMISTRY AND DAILY LIFE Chlorine is prepared by abstracting hydrogen from hydro- chloric acid. This may be done by simply passing the electric current through an aqueous hydrochloric acid solu- FIG. 12. A carboy of hydrochloric acid. tion, whereupon chlorine is evolved on one pole and hydro- gen on the other. The common way, however, is to oxidize hydrochloric acid by means of a suitable oxidizing agent. As such manganese dioxide is usually chosen, thus : Mn0 2 manganese dioxide 4HC1 = MnCl 2 hydrochloric acid manganous chloride H 2 O water C1 2 chlorine THE HALOGENS 55 The manganese dioxide and hydrochloric acid are gently heated together in a flask or retort. Chlorine is a greenish yellow gas with a very pungent odor. It is extremely active chemically, forming chlorides with many of the elements by direct union. So with sodium it forms sodium chloride, XaCl ; with copper, cupric chloride, CuCl 2 ; with phos- phorus, phosphorus chloride, PC1 3 ; with sulphur, sulphur chloride, S 2 C1 2 ; etc. On water chlorine reacts in the sunlight, forming oxygen and hydrochloric acid, thus : H 2 O + C1 2 = 2 HC1 + O The oxygen thus set free oxidizes many substances, destroy- ing colors, microorganisms, etc. For this reason chlorine water (a solution of chlorine gas in water) is a bleaching agent and an antiseptic. When chlorine gas is conducted into slaked lime, so-called chloride of lime, bleaching powder, is formed, thus : Ca(OH) 2 + C1 2 = CaCl(OCl) + H 2 O slaked lime chlorine chloride of lime water bleaching powder Bleaching powder is a powerful antiseptic and bleaching agent. Its action depends upon the fact that it will readily yield its oxygen and pass over into calcium chloride, CaCl 2 , thus : The oxygen destroys the color of many dyestuffs, and also kills microorganisms. Bleaching powder is manufactured on a large scale commercially, and is used in bleaching fabrics, paper pulp, etc. It is also used to rid the soil of undesirable organisms. Bromine is mainly found as sodium bromide, NaBr, in connection with sodium chloride. It may be prepared by methods that are entirely similar to those used in making 56 CHEMISTRY AND DAILY LIFE chlorine. Thus hydrobromic acid is formed when sulphuric acid acts on sodium bromide : NaBr + H 2 SO 4 = NaHSO 4 + HBr sodium sulphuric acid sodium hydrobromic bromide acid sulphate acid By treating manganese dioxide with hydrobromic acid, bromine is obtained : MnO 2 + 4 HBr = 2 H 2 O + MnBr 2 + Br 2 manganese hydrobromic water manganous bromine dioxide acid bromide Bromine forms bromides by direct union with many of the elements. Potassium bromide, KBr, is used in medicine and photography. Bromine itself is a brownish liquid of sp. gr. 3.188 at 0. It boils at 59 C. About 250,000 pounds of it are produced yearly in the United States. It is used as an antiseptic, also in making aniline dyes, bromides, etc. Hydrobromic acid when pure is a colorless gas soluble in water. Its properties are similar to those of hydrochloric acid, but it is weaker than the latter. Bromine is set free from bromides by chlorine, thus : NaBr + Cl NaCl + Br sodium chlorine sodium bromine bromide chloride Bromine turns starch paste yellow. Iodine is found in the ashes of sea weeds, also as sodium iodate, NaIO 3 , in connection with Chili saltpeter. Iodine is a crystalline, grayish black solid having almost a metallic luster. It dissolves sparingly in water, but copiously in alcohol. The alcoholic solution is called tincture of iodine. It is used in medicine as an antiseptic and counter irritant. With the metals and some of the other elements iodine forms iodides. Thus we have sodium iodide Nal, potassium iodide KI, silver iodide Agl, calcium iodide CaI 2 , phosphorus THE HALOGENS 57 iodide Pis, etc. From sodium iodide, iodine may readily be obtained by treatment with sulphuric acid. Thus hydro- gen iodide, hydriodic acid, HI, is formed, which, however, at once attacks the sulphuric acid present, forming sul- phorous acid, H 2 SO 3 , and water, thus : Nal + H 2 SO 4 = NaHSO 4 + HI sodium sulphuric acid sodium hydriodic iodide acid sulphate acid and H 2 SO 4 + 2 HI = H 2 S0 3 + H 2 O + 21 sulphuric hydriodic sulphurous water iodine acid acid acid Iodine turns starch paste blue. This fact is used in testing for starch. It is evident that one may also use starch paste in testing for iodine. The vapors of iodine are of a beautiful violet color, from which fact the element has received its name. The thyroid gland contains iodine in the form of a complex compound, thyroiodine. In the treatment of goiter and other diseases connected with the thyroid gland, iodine compounds, like potassium iodide, or extract of the thyroid gland of the sheep, are frequently prescribed by the physician. Iodine may be liberated from iodides by treating the latter with bromine or chlorine, thus : KI + Cl KC1 + I potassium chlorine potassium iodine iodide chloride QUESTIONS 1. Name the halogens, and one important compound of each found in nature. 2. (a) How may hydrofluoric acid be prepared ? (6) What use is made of this compound ? 3. Describe hydrochloric acid and tell how it is prepared. 4. What is hydrochloric acid used for? What other names does this compound have ? 58 CHEMISTRY AND DAILY LIFE 5. How may chlorine be made ? Describe this element. 6. Explain how chlorine bleaches. 7. What is bleaching powder ? Upon what fact does its action depend ? 8. What use is made of bromine ? How may it be prepared ? 9. Compare the properties of iodine with those of the other elements. 10. Mention some of the uses of iodine. Compute how much iodine would be obtained when 100 grams of chlorine act upon an excess of potassium iodide. CHAPTER VII SULPHUR, PHOSPHORUS, ARSENIC, ANTIMONY, AND BISMUTH Sulphur is found in nature in the free or uncombined state in large quantities. The largest deposits occur in Sicily, Texas, and Louisiana. Sulphur was well known even in ancient times. It frequently occurs around the brims of craters of volcanoes, and so is also called brimstone. In the market it may be obtained in sticks as roll sulphur, and also in powdered form as " flowers of sulphur." Sul- phur melts at 114. 5, forming a lemon-colored liquid, which on further heating turns brown and becomes viscous at about 160 to 200 C. Finally on further heating this liquid again becomes limpid and boils at 450 C., giving off brown vapors. Sulphur will burn in the air or in oxygen, forming sulphur dioxide, thus : g + Q 2 = S Q 2 Sulphur dioxide is a gas of suffocating odor. In it living beings cannot exist, and so it is used as a disinfecting agent. As such it is cheap. It is very commonly used in fumigating rooms and houses .that have been occupied by persons hav- ing Contagious diseases. FIG. 13: A can filled with liquid sul- phur dioxide. tor this purpose it may be had on the market in liquid form in tin cans, Fig. 13. These are soldered shut and have a small short lead tube attached to them, the end of which can be nipped off. The 59 60 CHEMISTRY AND DAILY LIFE gas can then make its escape into the room. The latter should be kept tightly closed for several hours after it has been filled with sulphur dioxide. The gas may also be pre- pared by burning sulphur in the rooms to be disinfected. Care must be taken in this process not to set the house on fire (see Chapter XXI). Sulphur dioxide is also used as a bleaching agent. It is especially used for bleaching straw, feathers, and fabrics FIG. 14. Pumping sulphur from a Louisiana sulphur well. SULPHUR, PHOSPHORUS, ARSENIC, ANTIMONY 61 that would be injured by bleaching powder. It used to be employed also in bleaching food products, such as dried apples, and grains discolored by fire, for example; but this is not to be recommended on account of the poisonous nature of the substance. Sulphur dioxide bleaches by uniting chemically with the substance whose color is altered. The action is thus quite different from that of bleaching with either hydrogen peroxide or bleaching powder, which bleach because of oxidation of the coloring matter. Sulphur is also used in making fireworks, black gunpowder, both soft and hard rubber, carbon bisulphide, lime sulphur mixtures for spraying shrubs and trees, and sulphuric acid. Thousands of tons of the latter compound are produced annually. In its production sulphur is first converted into sulphur dioxide by burning it in the air. The sulphur dioxide is then mixed with air and this mixture is passed over finely divided platinum heated from 400 to 450 C. In contact with this hot, finely divided platinum the sulphur dioxide unites further with the oxygen of the air, forming sulphur trioxide, thus : SO 2 + O = SO 3 Sulphur trioxide forms long, white, crystalline fibers, which melt at 14.8 C. It takes on water very greedily, forming sulphuric acid, thus : SO 3 + H 2 = H 2 SO 4 Sulphuric acid, also called oil of vitri61, is a heavy viscous liquid, being 1.838 times as heavy as water at 15 C. It has great affinity for water and will dissolve in the same with liberation of much heat. Care must consequently be used in pouring sulphuric acid into water. Always pour the acid 62 CHEMISTRY AND DAILY LIFE gradually into the water (never the water into the acid) and stir well after each addition of acid. If the water is poured into the acid, or if too much of the acid is poured into the water at a time, the heat evolved is so great that liquid is liable to be thrown out of the container and injure the per- son. Sulphuric acid is a very important article of com- merce. With most of the metals, ammonia, and other basic substances, it forms salts called sulphates. Thus we have copper sulphate CuS0 4 , calcium sulphate CaSO 4 , barium sulphate BaSO 4 , ammonium sulphate (NH 4 ) 2 SO 4 , etc. Be- sides its ability to unite with basic substances to form sulphates, the most striking property of sulphuric acid is its attraction for water, which has already been mentioned. So great is the affinity of sulphuric acid for water that it will abstract the latter from wood, sugar, starch, meat, and in fact from all other animal and vegetable substances, thus charring them. Sulphuric acid is further used in making super- phosphate fertilizers, in preparing hydrochloric acid and soda by the LeBlanc soda process, in manufacturing ether, aniline dyes, guncotton, dynamite, smokeless powder, and many other compounds in which a strong drying agent is required. Up- wards of four million tons of sulphuric acid are manufactured annually. Besides the process of making sulphuric acid, which has already been described and which is known as the contact process, sulphur dioxide is also oxidized by means of nitric acid and oxides of nitrogen in the presence of water vapor. This is an older process which is carried on in lead- lined chambers, lead being but very slightly attacked by sulphuric acid. The lead chamber process is still used to a large extent, especially for making sulphuric acid of 60 to 78 per cent strength. Sulphuric acid will react with many salts, forming sulphates and liberating the acids of such salts, thus : SULPHUR, PHOSPHORUS, ARSENIC, ANTIMONY 63 NaCl + H 2 S0 4 = NaHS0 4 + HC1 sodium chloride sulphuric acid acid sodium hydrochloric sulphate acid KN0 3 + H 2 SO 4 = KHS0 4 + HNO 3 potassium nitrate potassium acid nitric acid sulphate Fuming sulphuric acid, also called Nordhausen sulphuric acid, is sulphuric acid in which additional sulphur trioxide has been dissolved. Sulphurous acid is formed when sulphur dioxide dissolves in water, thus : S Q 2 + HaO = jj^ From it salts called sulphites are derived. Thus we have sodium sulphite Na 2 SO 3 , sodium acid sulphite, also called sodium bisulphite, NaHS0 3 , calcium sulphite CaSO 3 , etc. The sulphites are used in photography and in the preparation of wood pulp in the paper industry. Sulphites also used to be employed as food preservatives, but this is now for- bidden by law, for they are injurious to health. Hydrogen sulphide is a compound consisting of hydro- gen and sulphur. Its composition is expressed by the formula H 2 S. This gas is formed in rotten eggs, in manure pits, and wherever animal and vegetable matter is decaying without proper access of air. Hydrogen sulphide is poi- sonous. It dissolves in water, one volume of the latter absorbing about three volumes of the gas. Hydrogen sul- phide is 1.19 times as heavy as air. It will burn, forming water and sulphur dioxide, thus : H 2 S + 3 O = H 2 + S0 2 Hydrogen sulphide may be prepared by treating ferrous sulphide with hydrochloric or sulphuric acid, thus : FeS + 2 HC1 = FeCl 2 + H 2 S FeS + H 2 S0 4 = FeS0 4 + H 2 S 64 CHEMISTRY AND DAILY LIFE When introduced into solutions of salts of the heavy metals, hydrogen sulphide forms characteristic precipitates consisting of the sulphides of the metals, and it is consequently much used in testing for heavy metals in chemical analysis, thus : CuS0 4 + H 2 S = CuS + H 2 S0 4 copper sulphide 2 AsCl 3 + 3 H 2 S = As& + 6 HC1 arsenious sulphide Carbon bisulphide is a colorless liquid of specific gravity 1.26. It boils at 47 and is very inflammable. Its vapors mixed with air are explosive, and consequently no flame or spark must be present when carbon bisulphide is being used. It will riot dissolve in water. It is used as a solvent for fats and oils, also for the extermination of ants and other insect pests. Carbon bisulphide is prepared by heating carbon and sulphur together out of contact of the air, usually in an electric furnace. The gas formed is then condensed by means of cold water. Sulphur occurs in nature also as calcium sulphate CaSO 4 , gypsum CaS0 4 . 2 H 2 O, and also in combination with metals as sulphides like iron pyrites or fool's gold FeS 2 , sulphide of lead or galenite PbS, zinc sulphide or black jack ZnS, etc. From pyrites sulphur dioxide is obtained by heating the mineral in the air. Much of the sulphur dioxide used in the manufacture of sulphuric acid is obtained from this source. Nearly two and one half million tons of gypsum are pro- duced annually in the United States alone. It is used as land plaster, as wall plaster, stucco, etc. Finally all plants and animals contain sulphur. With- out it they could not live. Sulphur is a constituent of albumen. Thus muscular tissues, horns, hoofs, hair, nerves, eggs, and seeds contain sulphur. In urine sulphur is present in the SULPHUR, PHOSPHORUS, ARSENIC, ANTIMONY 65 form of sulphates, for as the animal lives, the sulphur com- pounds in its tissues are continually oxidized to sulphates, which are then eliminated by means of the kidneys. Some plants, like onions, garlic, mustard, skunk cabbages, radishes, contain odoriferous sulphur compounds. In order to thrive well they require a soil which contains an adequate amount FIG. 15. A group of sulphur crystals on limestone. of sulphur. When plants and animals decay, the, sulphur compounds they contain gradually decompose, forming hydrogen sulphide when the sulphur is relatively abun- dant and access to the air is limited, and sulphates when there is proper access of air. When these sulphates are again put into the soil, as in manuring or in treating with land plaster (i.e. gypsum), the rootlets of plants again take them up, and transform them into the various sulphur F 66 CHEMISTRY AND DAILY LIFE compounds of their tissues, and thus the sulphur cycle is completed. Sulphur springs contain hydrogen sulphide, which is easily detected by its foul odor. Many mineral waters contain calcium sulphate, some contain sodium sulphate, Na 2 SO 4 , others like those at Epsom contain magnesium sulphate, MgSO4. The latter are bitter to the taste and have a laxative effect. Magnesium sulphate is called Epsom salt, and is sometimes taken as a purgative. Phosphorus never occurs in nature in the free state. It is, however, quite widely distributed in small quantities in compounds. It occurs as calcium phosphate, Ca 3 (PO 4 )2, in phosphate rock in the Carolinas and some of the Western states. Small amounts of phosphates occur also in all fertile soils, in iron ores (from which it passes into blast furnace slags when the ores are re- duced), and in many granitic rocks, clays, etc. Bones consist of calcium phosphate to the extent of about 80 per cent. Phos- phorus is an important and indispensable part of the tissues of all plants and animals. In the nerves, brain, blood, muscles, and in fact in all animal parts that are concerned in locomotion or reproduction, phosphorus is found. In seeds phosphorus is always present, particularly in the embryos, and without phosphorus plants cannot grow, hence FIG. 16. A piece of phosphate rock. SULPHUR, PHOSPHORUS, ARSENIC, ANTIMONY 67 its importance in the soil. In the bodies of plants and animals phosphorus is combined with carbon, hydrogen, oxygen, nitrogen, and sulphur. The nerves and the brain are especially rich in a complex compound, leci- thine, whose composition is expressed by the formula FIG. 17. Barley growing with and without proper food. ( 1 ) Complete manure. (4) No potassium. (2) No nitrogen. (5) No calcium. (3) No phosphorus. (6) No magnesium. 68 CHEMISTRY AND DAILY LIFE . As animals live these complex compounds are gradually oxidized, phosphates being formed, which are then eliminated by the kidneys. Thus urine always contains phosphates, especially calcium and potassium phosphates. When these are then returned to the land as fertilizer, they are again taken up by the roots of plants and converted into complex phosphorus compounds in their tissues and seeds. These in turn are eaten by animals, and so the phosphorus cycle com- pletes itself. It is clear, however, that considerable amounts of phosphorus are never returned to the land, being washed away continually together with other valuable fertilizer material into the waterways, which finally empty into the ocean. The sewage of our large cities in particular represents a colossal waste. In order to keep up the supply of phosphorus, bones, blast furnace slags, and ground phosphate rock are employed as fertilizers. When bones are treated with sul- phuric acid, so-called superphosphate is produced, thus : Ca 3 (P0 4 ) 2 + 2H 2 SO 4 = 2 CaSO 4 + CaH 4 (PO 4 ) 2 calcium phosphate sulphuric acid superphosphate This is more soluble in the waters of the soil and hence more available to plants. Phosphorus itself is produced by heating calcium phosphate with silica (sand) and carbon. These are finely ground and intimately mixed and then heated out of contact with the air in earthenware retorts or in a suitable electric furnace. The following change takes place : Ca 3 (PO 4 ) 2 + 3 SiO 2 + 5 C = 3 CaSiO 3 + 5 CO + 2 P calcium silica carbon calcium carbon phos- phosphate silicate monoxide phorus Phosphorus is a yellowish, translucent, waxlike solid. Under water it may be melted at 44 C. It catches fire in the air at about 35 C., and hence it is always kept and trans- SULPHUR, PHOSPHORUS, ARSENIC, ANTIMONY 69 ported out of contact with the air, usually under water in- air-tight cans. Phosphorus is very poisonous, 0.1 gram being sufficient to produce death in the case of an adult. Phos- phorus burns are dangerous, and they heal very slowly. Phosphorus should consequently be handled with great care. Forceps should always be used when handling phos- phorus, and small amounts, usually not larger than a pea, should be employed in experiments by the beginner. Phosphorus is used to poison rats and other vermin, and in making matches. For the latter purpose yellow phos- phorus should not be employed, for the workers In match factories are frequently poisoned by working with it, the poisoning being evidenced by an enlargement of the liver and a gradual destruction of the jawbones. When heated out of contact with the air to from 250 to 300 C., yellow phosphorus changes to a red powder, called red phosphorus. This allotropic form of phosphorus is far less dangerous than the yellow variety. It will not take fire in the air till heated to about 200 C., and moreover it is not poisonous. It is frequently employed in the chemical laboratory in preparing phosphorus compounds, and it is also used in making safety matches. The latter commonly contain a mixture of potas- sium chlorate, potassium bichromate, powdered glass, and glue or dextrine, as the match head, whereas on the friction surface on the box there is a mixture of red phosphorus, antimony sulphide, manganese dioxide, and glue. The powdered glass increases the friction, which raises the tem- perature to the point at which the match head bursts into flame. Safety matches ignite only by rubbing on the specially prepared surface on the box. They were first prepared in Sweden and so are often called Swedish matches. Matches containing yellow phosphorus in the head will ignite by rubbing on any surface. They are dangerous because they may cause 70 CHEMISTRY AND DAILY LIFE conflagrations, and also because of their poisonous nature many lives are sacrificed in manufacturing them. Phos- phorus matches first came into use in 1832, and now some 3000 tons of phosphorus are annually used up in making matches. Phosphoric acid is produced by treating calcium phosphate with sulphuric acid, thus : Ca 3 (PO 4 )2 + 3H 2 SO 4 = 3CaSO 4 + 2 H 3 PO 4 calcium phosphate sulphuric acid calcium sulphate phosphoric acid It is a sirupy liquid of specific gravity 1.88. It is a pro- nounced acid and forms phosphates with the metals, with ammonia, and other bases. On heating phosphoric acid it loses water and forms metaphosphoric acid, HPO 3 , thus : H 3 P0 4 = HPO 3 + H 2 O Metaphosphoric acid is analogous to nitric acid, HNO 3 . When phosphorus is burned in the air or in oxygen, phos- phorus pentoxide P 2 O 5 is formed, thus : 2 P + 5 O = P 2 O 5 The latter on treatment with water forms phosphoric acid, thus: P 2 5 +-3H 2 = 2H 3 P0 4 As phosphorus is a valuable constituent of fertilizers, the percentage content in the latter should always be considered in making purchases. The per cent of phosphorus is some- times stated as per cent P, which means the element itself, but more frequently it is stated on the label as per cent phosphorus pentoxide, P 2 O 5 , which is erroneously called phosphoric acid. It is to be remembered that P 2 O 5 contains only 62 pounds of phosphorus to 80 pounds of oxygen. That is to say, it contains only 62 pounds of phosphorus in SULPHUR, PHOSPHORUS, ARSENIC, ANTIMONY 71 every 142 pounds. Dealers like to express the phosphorus content of fertilizers as per cent of P2O 5 , for, of course, the figure appears much larger than when expressed as per cent P. Arsenic is a steel-gray to black, brittle substance having a metallic luster. It burns in the air emitting a garlic-like odor and thus forms a white powder, arsenious oxide, or white arsenic, As 2 O 3 , which floats in the air as a white cloud. Indeed, white arsenic is the most important compound of arsenic. It is quite common on the market and relatively inexpensive. Like all compounds of arsenic, it is very poi- sonous. White arsenic, arsenious acid, or " arsenic," as it is also often called in commerce, is used as rat poison. About 1500 tons are produced yearly in the United States. It is slightly soluble in water and has a sweetish, sickening taste. The antidote for arsenic poisoning is freshly precipitated ferric hydroxide. This forms an insoluble mass with ar- senious oxide. Arsenic compounds are used almost entirely as poisons for vermin and insect pests. Two compounds of arsenic are at present very commonly employed for this purpose, namely, Paris green, whose composition is Cu 3 As 2 O 6 . Cu(C 2 H 3 O 2 )2, being a double salt of cupric arsenite and cupric acetate, and lead arsenate Pb 3 (AsO 4 ) 2 . Both Paris green and lead arsenate are but sparingly soluble in water. They are used in exterminating potato bugs and various other insects. Directions for preparing the proper mixtures for such spraying will be given later. It should be kept in mind that both compounds are strong poisons and must be used with proper care. Antimony is a brittle metal. It is used in making type metal and Babbitt metal for machine bearings. As molten antimony solidifies, it expands and fills the molds, hence it is valuable in many alloys ; moreover, it is hard and when 72 CHEMISTRY AND DAILY LIFE mixed with lead and tin forms an alloy which is particularly suitable for type, the composition being 25 per cent anti- mony, 50 per cent lead, and 25 per cent tin. Babbitt metal contains 70 to 90 per cent lead, alloyed with tin and anti- mony. Bullets and shot consist of lead alloyed with from 0.2 to 0.4 per cent arsenic. Bismuth, too, is a brittle metal. It is used in preparing low-melting alloys which are employed in making readily fusible plugs in the pipes of automatic sprinklers for fire protection. So, for instance, Wood's metal, consisting of 1 part tin, 2 parts lead, 1 part cadmium, and 4 parts bismuth, melts at 60. 5 C. Bismuth compounds are used in medi- cine, the most common preparation being bismuth sub- nitrate, BiO . NO 3 . BiO . OH, a white powder which is diffi- cultly soluble in water and practically tasteless. It is pre- scribed in cases of dysentery and other disturbances of the alimentary canal. It is also sometimes used as a face powder. QUESTIONS 1. How does sulphur occur in nature ? 2. Describe how this element behaves on heating. 3. What is formed when sulphur is burned in the air ? Describe the properties and uses of this substance. 4. What use is made of sulphur ? 6. What are the two methods by which sulphuric acid is manu- factured on a large scale ? 6. Describe the properties of sulphuric acid. What use is made of this substance ? 7. What are sulphites? Sulphates? Give an example of each. 8. What is hydrogen sulphide ? How is it prepared ? 9. How is carbon bisulphide made ? What is it used for ? 10. Discuss the role of sulphur in plants and animals. SULPHUR, PHOSPHORUS, ARSENIC, ANTIMONY 73 11. In what forms is phosphorus found in nature ? 12. Discuss the role of phosphorus in plant and animal life. 13. What use is made of phosphorus in the arts ? 14. What is superphosphate fertilizer ? 15. How much phosphoric acid, HsPO-i, could be produced from one ton of calcium phosphate, Ca3(P0 4 )2, the essential compound in phosphate rock ? 16. How much phosphorus is there in 1 pound of phosphorus pentoxide, P20s? How much phosphoric acid, H 3 P04, could be made from 1 pound of P20 5 ? 17. What is white arsenic ? 18. In what other commercial compounds does arsenic occur? What are these used for? 19. How much arsenic is there in a pound of Paris green ? 20. What are antimony and bismuth, and what use is made of them? CHAPTER VIII BORON AND SILICON Boron occurs in nature in boric acid and its sodium and calcium salts. It is never found in the uncombined state. The element boron is a brown, brittle solid. It may be dis- solved in molten aluminum, and when the mass cools boron separates out in form of crystals that are almost as hard as diamond. Boric acid has the composition H 3 BO 3 . It may be prepared from borax, Na 2 B 4 O 7 . 10 H 2 O, by treating a hot concentrated solution of the latter with either hydrochloric or sulphuric acid. The boric acid separates out in the form of crystalline flakes on cooling, thus : Na 2 B 4 7 + 2 HC1 + 5 H 2 O = 2 NaCl + 4 H 3 BO 3 sodium hydrochloric water sodium boric acid tretraborate acid chloride Boric acid and borax are by far the most important compounds of boron. Boric acid is known also as boracic acid. It dissolves in water. At room temperature the saturated solution contains about 4 per cent. It is frequently used as an antiseptic wash, especially in treating the eyes, for it is a very mild agent. Compresses of saturated boric acid solu- tion are frequently applied on parts of the body that are swollen because of blood poisoning. Powdered boric acid feels slip- pery to the touch, and consequently is at times used to spread on the floors of dance halls. It has also been used to preserve milk, meats, fish, etc. ; but as it is injurious to health, the practice is now forbidden by law. Boric acid has prac- 74 BORON AND SILICON 75 tically no effect on litmus, for it is an extremely weak acid ; it turns turmeric paper reddish brown, however, and this serves as a delicate test for boric acid. When the turmeric paper thus reddened is afterward treated with caustic alkali solution, it turns to a dark greenish black. Borax is alkaline toward litmus, for it is a salt of a very weak acid and a powerful base. The salt forms beautiful crystals, which crumble to pow- der on standing in the air, be- cause they lose water. Over forty thousand tons of borax are produced in the United States annually. The chief deposits of colemanite, Ca 2 B 2 On . 5 H 2 O, from which much borax is made, occur in California and Oregon. In the laundry, borax is used for softening water and for en- hancing the gloss of starch in ironing. It is also used as an antiseptic, as a mordant in dyeing fabrics, and as a flux in welding iron and brazing metals together. Like boric acid, it has also been used for preserving foods, but this practice is to be strongly condemned. In making glazes and enamels for pottery and enameled iron ware, borax and boric acid are often used, also in the production of certain kinds of hard glass. While boron compounds after all are found in relatively small quantities on the earth, compounds of silicon are ex- tremely abundant, for they make up by far the larger portion FIG. 18. A group of borax crystals. 76 CHEMISTRY AND DAILY LIFE of the solid part of the earth. The element silicon never occurs in the uncombined form. It is a brittle solid having metallic luster. Its specific gfavity is 2.49. It is so hard that one can scratch glass with it. It is now made on a large scale by heating silicon dioxide, quartz sand, with coke in the electric furnace, thus : SiO 2 + 2C 2 CO + Si silicon dioxide carbon carbon monoxide silicon At present silicon is used mainly in the manufacture of steel, where it serves as a reducing agent. FIG. 19. Molten silicon as it is being tapped from an electric furnace. Silicon dioxide, SiO 2 , also called silica, is the most impor- tant compound of silicon. Quartz, quartzite, flint, chalcedony, opal, amethyst, agate, carnelian, sandstone, white sand, espe- cially that on the seashore and the desert, are almost entirely BORON AND SILICON 77 silicon dioxide. Silica is hard and brittle, and is much used as an abrasive. Sandpaper consists of quartz sand glued on paper. Silica may be melted to a clear glass in the oxyhy- drogen flame. Utensils of such glass are still quite costly. They are mainly used in laboratories. Although quartz glass is quite brittle, it will not break when exposed to sudden changes of temperature. So a dish of quartz glass may be made red hot and then thrust into cold water without break- ing the dish. Quartz sand is also employed in making glass, porcelain, portland cement, and ordinary lime mortar. FIG. 20. Quartz crystals from Hot Springs, Arkansas. When pulverized quartz sand is fused with soda (sodium carbonate), sodium silicate is formed, thus: SiO 2 + Na 2 CO 3 = Na 2 SiO 3 + CO 2 silica soda sodium silicate carbon dioxide water glass Sodium silicate is also called water glass. It is readily soluble in water, and its solutions are employed in preserving eggs, and in cementing asbestos fibers and asbestos paper together. When sodium silicate solution is treated with 78 CHEMISTRY AND DAILY LIFE hydrochloric or some other strong acid, silicic acid is formed, and this separates from concentrated solutions in the form of flakes or jelly, thus : Na 2 Si0 3 + 2HC1 = H 2 SiO 3 sodium silicate hydrochloric acid silicic acid + 2 NaCl sodium chloride At ordinary temperatures silicic acid is a very weak acid. It does not affect litmus, and has no taste. In water it is but very slightly soluble. However, when the sodium silicate solution used to prepare the silicic acid (see above equation) is only 1 per cent strong, silicic acid will not separate on acidifying with hydrochloric acid. By placing this solution in an animal bladder, like a hog's bladder, and immersing this in a pail of water, the sodium chloride and excess of hydrochloric acid in the solution will pass through the walls of the bladder into the water on the outside of the latter. By renewing the water in the pail constantly, there will finally remain behind in the bladder only silicic acid. In- stead of a bladder, a bag made of parchment paper could also be used. This process of thus separating the silicic acid from the other ingredients in the solution was first used by Thomas Graham. It is called dialysis. By carefully evapo- rating off some of the water, the silicic acid solution may be concentrated till it contains about 10 per cent silicic acid. However, the latter is very apt to separate in the form of a jelly on standing. Salts of silicic acid are exceedingly abundant in nature. FIG. 21. Preparing silicic acid by dialysis. BORON AND SILICON 79 The only soluble salts of this acid are those of sodium and potassium and these are prepared artificially. The naturally occurring silicates are insoluble. Among some of the most important of these may be mentioned : potash feldspar KAlSi 3 O 8 , soda feldspar, NaAlSi 3 O 8 , lime feldspar CaAl 2 Si 2 O 8 , kaolin or day H 2 Al 2 Si2O8 . 2 H 2 O, serpentine Mg 3 Si 2 O 7 , oliv- ine Mg 2 SiO 4 , meerschaum Mg 2 Si 3 O 8 . 4 H 2 O, soapstone or talc Mg 3 H 2 Si 4 Oi 2 , asbestos Mg 3 Si 2 O 7 . 2 H 2 O, hornblende Mg 2 CaFeSi 4 Oi 2 , and mica KH 2 Al 3 Si 3 Oi 2 . Granitic rocks con- sist of crystals of quartz, feldspar, and mica. Our fine clay soils are formed by the disintegration of the granitic rocks, partly by the action of glaciation in earlier geological ages, partly by the process of weathering, that is, by the gradual action of water, air, and wind, particularly by alternate freezing and thawing of water in the crevices of the rocks. Because of their content of potassium and also of phos- phorus, for granitic rocks and the clay that have come from their disintegration always contain minor amounts of cal- cium phosphate, clay soils are in general quite desirable for agricultural purposes. While sand and the naturally occurring silicates are all only very slightly soluble in water, yet they are by no means entirely insoluble. So all terrestrial waters like those of springs, wells, rivers, lakes, and the ocean, as well as all soil waters, contain silicates in solution. From these dissolved silicates, plants derive in a considerable part the mineral matter which constitutes the ash that remains when they are in- cinerated. Silica itself is frequently met in the ash of plants. It is found particularly in the ash of the stalks of the cereals, of grasses, of bamboo and other canes, to which it gives stability. Often half of the ash consists of silica. In feathers, in the hair of animals, and in the shells of various crustaceans, silica abounds. About 40 per cent of the ash 80 CHEMISTRY AND DAILY LIFE of the feathers of birds consists of silica. Deposits of dia- tomic or infusorial earth consist of the siliceous remains of minute organisms. This infusorial earth is used in making dynamite, which is essentially nitroglycerine mixed with infusorial earth and then molded into sticks. In the plant and animal bodies silica is doubtless combined with other elements in the form of quite complex compounds. The exact nature of these has not yet been determined. Hydrofluoric acid will attack silica, forming a volatile compound, silicon tetrafluoride and water, thus : SiO 2 + 4HF SiF 4 + 2 H 2 silica hydrofluoric acid silicon tetrafluoride water For this reason hydrofluoric acid is much used in the lab- oratory in analyzing silicates. It is also employed in etching silica, silicates, glass, and porcelain. The etching depends upon the fact that silicon tetrafluoride is formed, and that fluorides of any bases that may be present are simultaneously produced. The art of thus etching glass has been known for centuries. QUESTIONS 1. What is borax ? What use is made of it ? 2. How is boric acid prepared ? What is it used for? 3. What is the most abundant compound of silicon? How much silicon do 100 Ib. of it contain ? 4. Make a list of the most important things in the manufacture of which quartz is used. 5. What is water glass ? How is it made, and for what purpose is it used ? 6. How are clay soils formed? Why are they desirable for agricultural purposes ? 7. Discuss the occurrence of silica in natural waters, also in plants and animals. 8. What is the action of hydrofluoric acid upon silicates ? CHAPTER IX CARBON AND ITS, COMPOUNDS WHEN a porcelain plate is held in the flame of a candle, carbon deposits on the plate in the form of soot, which is also called lampblack. When in the process of baking pota- toes, apples, bread, or meats they are left too long in a some- what overheated oven, only charred masses which are largely carbon remain. In fact, all animal or vegetable matter when similarly heated yields residues of carbon, showing that the latter element is an essential constituent of all living beings. Such charred remains of animal or vegetable matter always weigh much less than the original material, which is due to the fact that a large portion of the latter has been volatilized during the process of heating. What remains is very largely carbon plus the mineral constituents, i.e the ash. On continued heating in the air the carbon in the charred mass gradually burns off, and there is finally left nothing but the ash. By heating wood out of contact with the air in ovens charcoal is pro- duced. When coal is similarly heated, coke is formed; and when bones are thus treated, bone black results. In like manner blood charcoal may be made from blood. Charcoal is generally G 81 FIG. 22. Making charcoal. A vertical section through the center of a pile. 82 CHEMISTRY AND DAILY LIFE produced by piling up wood in a suitable manner, covering the whole with earth to exclude undue access of air, and then setting fire to the wood. Thus a portion of it burns, and the rest is merely charred because it is heated practi- cally out of contact with the air. Charcoal, coke, and bone black ahcays contain, besides carbon, the mineral substances, i.e. the ash, of the original material that ivas charred. FIG. 23. Charcoal burning in Bavaria. Diamond is crystalline carbon. It is the hardest substance known. When pure it is colorless and has a high index of refraction, for which reason it is greatly prized as a gem. Less desirable pieces of diamond are used for making drills to drill rocks. Graphite also occurs in nature and is not infrequently in a fairly pure state. It is soft, black, and " soapy " to the touch. It is used as a lubricant, also for making " lead " pencils. Graphite is also termed plumbago, CARBON AND ITS COMPOUNDS 83 which name originated because it was formerly thought that the substance contained lead. Graphite is now made arti- ficially by heating carbon very highly out of contact with the air in the electric furnace, and then allowing it to cool very slowly. Under these conditions the amorphous carbon is changed to graphite. The latter may now be obtained in pieces of almost any size desired. These can readily be worked into any form required by means of lathes, milling machines, or bench tools. Much of this artificial graphite is now used in making pencils, electrodes for electrolytic work, resistances, etc. Charcoal will absorb gases, and hence it is frequently used to take odoriferous gases out of cisterns, pits, vaults, etc. The cause of such gases ought in all cases to be removed, however, if possible. To hang a bag full of charcoal in an ill-smell- ing cistern, when the latter contains decaying organic matter that ought to be taken out, is obviously not the proper way of dealing with the problem. In the sick room charcoal dressings are at times employed to absorb fetid odors issuing from ulcers. Such dressings need frequent renewal to be efficient. Bone black consists of only about 8 to 12 per cent carbon, the remainder being largely calcium phosphate. Bone black absorbs many coloring matters from solutions, and hence is used in sugar refining to remove the brown color from the sugar solutions. Blood charcoal is more expensive than bone black. It absorbs coloring matter better, and conse- quently is often employed in the chemical laboratory in experimental work. Like charcoal, bone black and blood charcoal also absorb odors. Coal represents the fossil remains of plants of the carbon- iferous and other geological ages. By the gradual loss of water and other volatile matter, like marsh gas, hydrogen, 84 CHEMISTRY AND DAILY LIFE etc., these vegetable remains have gradually been trans- formed to coal. Wood fiber, peat, brown coal, soft coal or bituminous coal, and anthracite or hard coal represent a series of gradations in the gradual process which has taken place in nature in the production of coal. Beginning with wood fiber, by gradual loss of oxygen, hydrogen, and some carbon as well, the other products just enumerated may be regarded as having been formed. So anthracite generally consists of about 95 per cent carbon, the rest being ash plus a few per cent of volatile matter. In soft coal there generally is about 80 per cent carbon and a very much higher amount of vola- tile matter than in hard coal. Soft coal is consequently used for making coal gas for heating and illuminating purposes. In this process the coal is heated out of contact with the air in iron retorts. In the latter finally coke remains, which also contains the mineral content or ash of the coal. The follow- ing represents the composition of a fairly typical sample of coal gas : methane or marsh gas, 34.5 per cent ; hydrogen, 49 per cent; carbon monoxide, 7.2 per cent; nitrogen, 2.3 per cent; oxygen, nil; carbon dioxide, 1.1 per cent; illumi- nants, benzene, etc., 5 per cent. Ammonia is also always formed in the manufacture of coal gas, as already stated in Chapter IV, but it is washed out of the gas before the latter is stored and admitted to the gas mains. All of the ammonia and ammonium salts of commerce are obtained from the gas liquors. When coal gas burns from an ordinary jet, a luminous flame results which deposits soot on a plate that is held in it. The luminosity of the flame is caused by the incandescent particles of carbon that are present in the flame. A hydrogen flame is colorless, but can be made luminous for a time by carefully blowing fine particles of soot into it. When suffi- cient oxygen is supplied to a coal gas flame, its luminosity CARBON AND ITS COMPOUNDS 85 disappears because the carbon in the flame is all consumed ; for the latter reason such a flame is hotter than the lumi- nous one. The Bunsen burner, Fig. 24, is an arrangement for burning gas by first mixing it with a sufficient amount of D FIG. 24. The Bunsen burner and its flame. air so that more perfect combustion is secured. The gas flows from the small orifice and air enters at the opening A. In the tube T both air and gas mix well, and this mixture is then ignited at B, the upper end of T. The gas burns with a flame that has an inner blue cone C and an outer zone D which is non-luminous. The inner zone contains 86 CHEMISTRY AND DAILY LIFE unconsumed gas, while in the outer zone practically com- plete combustion takes place. This is the hottest part of the flame. All gas flames intended for heating purposes are arranged on this principle. So, for example, the burners in the laboratories, the gas stoves and ranges, the gasoline heaters, etc., are all arranged so that the gas is first well mixed ivith air and then the mixture is burned from a jet. In the case of blast lamps air is blown into the gas by pressure from some source. A flame results only when a gas burns. Charcoal or coke never burn with a flame, for in their production all gases were expelled from them. Charcoal or coke in burning become red hot and glow till they are consumed so that only ash remains. When carbon burns in the air or in oxygen in excess, carbon dioxide results, thus : c + o 2 = co 2 carbon oxygen carbon dioxide When, however, the amount of oxygen is limited and an excess of carbon is present, carbon monoxide is formed, thus : C + O = CO carbon oxygen carbon monoxide Carbon monoxide is a gas which burns with a blue flame, yielding carbon dioxide, thus : CO + O = CO 2 carbon monoxide oxygen carbon dioxide Carbon monoxide is poisonous. When breathed it unites with the hemoglobin of the blood and produces death. Carbon monoxide is the most dangerous ingredient of coal gas. Its flame is quite hot. In " water gas," which is a mixture of CARBON AND ITS COMPOUNDS 87 hydrogen and carbon monoxide produced by passing steam over white-hot coke, thus : C + H 2 O = CO + H 2 , carton water water gas we have then a gas which is excellent for heating, but not at all suitable for illuminating purposes. By " enriching " this gas with benzene, or gas pro- duced from petroleum oils, it may be used as illuminating gas. With the Welsbach mantle lamp water gas may be used very well, for here a non-luminous Bunsen flame is employed, over which the mantle, consisting of a net- work of 1 per cent cerium oxide and 99 per cent thorium oxide, is placed. This " stocking " is heated to incandescence and emits a brilliant light. Producer gas consists of 28 to 30 per cent carbon monoxide, 63 per cent ni- trogen, and minor amounts of carbon dioxide. It is made by passing air over red-hot coke. The gas is readily made and is FIG. 25. A Welsbach mantle. frequently used as a source of heat in industrial processes. The nitrogen it contains is, of course, a drawback, for it acts simply to dilute the carbon monoxide. Carbon dioxide, as already stated, is contained in the atmosphere, and from this source the carbon that is found in living beings is really obtained. In the green leaf of the plant in the sunlight carbon dioxide from the air and moisture 88 CHEMISTRY AND DAILY LIFE react with each other, forming starch and eliminating oxygen, thus : 6 C0 2 + 5 H 2 = 6 2 + C 6 H 10 5 carbon dioxide water oxygen starch Starch is an article of food. It is present in large quantities in all cereals, potatoes, and other vegetables. As these are eaten the starch is changed to sugar in the alimentary canal, passes through the walls of the latter into the blood, and is utilized in building up tissues of the body. The latter as we live are again continually slowly oxidized by the oxygen which passes through the walls of the lungs into the blood as we breathe. Thus carbon dioxide is formed and this is exhaled with every breath. In this way carbon is returned to the air in the form of carbon dioxide and the carbon cycle has been completed. The plant can now again take up the carbon dioxide and transform it into starch. The latter can again be eaten, assimilated, oxidized to carbon dioxide in the body, and exhaled, and so on. The energy that keeps this process going obviously comes from the sun without whose light and warmth the plant could not live and continue its work of transforming carbon dioxide and water into starch and oxygen. Carbonated waters are waters charged with carbon dioxide, the latter gas being placed under pressure over water in a closed tank. When such pressure is released from the water, a considerable amount of the gas escapes in bubbles. Nat- ural springs whose waters are charged with carbon dioxide occur in nature. So, for example, at Colorado Springs the waters are highly carbonated. Soda water is water charged with carbon dioxide. It is so called because the gas employed was formerly made by the action of an acid on baking soda, thus : 2 NaHCO 3 + H 2 S0 4 = Na 2 SO 4 + H 2 O + 2 CO 2 baking soda sulphuric acid sodium sulphate water carbon dioxide CARBON AND ITS COMPOUNDS 89 Carbonated waters are refreshing to the taste and are conse- quently rightfully highly esteemed. Carbon dioxide from the fermentation industries is now commonly stored in steel cyl- inders and used for making carbonated beverages of all kinds. Since carbon dioxide occurs in the air and since water dissolves carbon dioxide, the latter is present in all natural waters. Carbonates are salts of carbonic acid. Carbon dioxide is really carbonic acid anhydride, that is, carbonic acid minus water. Carbonic acid is H 2 CO 3 , but it has never been isolated, for it decomposes readily into carbon dioxide and water, thus : H 2 C0 3 = C0 2 + H 2 carbonic acid carbon dioxide water The salts of carbonic acid are, however, quite common and stable under ordinary conditions. They are also of great im- portance. So limestone and marble consist of almost pure calcium carbonate, CaCO 3 . Magnesium limestone or dolo- mite consists of magnesium and calcium carbonate, MgCOs -f- CaCO 3 . These are valuable as building materials, (1) as building stones and (2) for making building lime. In the process of making lime the limestone is heated in a kiln out of contact with the air. Thus carbon dioxide is expelled, and lime, calcium oxide, CaO, remains : CaCO 3 on heating = CaO -f CO 2 calcium carbonate calcium oxide carbon dioxide limestone lime In slaking the lime, as for building purposes, making white- wash for walls, or producing liquids for spraying trees, etc. calcium hydroxide results and considerable amounts of heat are evolved as the chemical union of lime and water proceeds, thus: CaO + H 2 = Ca(OH) 2 calcium oxide water calcium hydroxide lime or slaked lime 90 CHEMISTRY AND DAILY LIFE When this slaked lime is mixed with sand, mortar is obtained, whose hardening or " setting " proceeds because water dries out FIG. 26. A homemade lime kiln. and at the same time the carbon dioxide of the air again acts on the slaked lime, forming calcium carbonate again, thus : Ca(OH) 2 + C0 2 = CaC0 3 + H 2 O calcium hydroxide carbon dioxide calcium carbonate water slaked lime The crystals of calcium carbonate tightly hold the grains of sharp sand in place, thus forming a solid mass. It is' clear that in order to set well, lime mortar must haw plenty of air so that it may secure the necessary carbon dioxide therefrom. Moreover, the air should be fairly dry in order that the water may dry out of the mortar. Plastering and bricklaying in cold, damp weather does not produce the best results. CARBON AND ITS COMPOUNDS 91 In many sands and soils grains of limestone, calcium car- bonate, are found. The shells of oysters, clams, and other shell- fish consist largely of calcium carbonate. Indeed, it is quite probable that the large beds of limestone and marble (which is merely a purer form of limestone) have been formed from the remains of mollusks, whose shells have been deposited from the sea in deep layers which have later, under heat and pressure, been transformed to the crystalline state. Carbon is further found in nature in petroleum, a liquid which consists of a mixture of compounds of carbon and hydro- gen. Compounds of the latter elements are called hydro- carbons. It is not easy to separate from petroleum in pure form any one of the mixture of hydrocarbons of which it consists. However, it is not at all difficult to separate petro- leum into groups of hydrocarbons whose boiling points are fairly close together. This is actually done in practice by the process of fractional distillation. Petroleum is distilled, and separate portions passing over between certain temperatures are collected in different receivers. In this way there are obtained from petroleum : rhigolene, which boils at about room tem- perature ; petroleum ether, boiling between 50 and 60 C. ; gasoline, boiling between 70 and 90 C. ; naphtha, boiling from about 90 to 120 C. ; benzine, boiling from 110 to 140 C. ; and kerosene, boiling between 150 and 300 C. Above the latter temperature heavy oil passes over as the distilla- tion proceeds. This is used for lubricating purposes. After the lubricating oils have been distilled off, vaseline passes over, and finally paraffine is prepared from the residue that remains in the retort. The oils prepared from petroleum are often called mineral oils to distinguish them from the fats and oils of plant or animal origin. The fats and oils from the latter sources are not hydrocarbons. They are compounds of car- bon, hydrogen, and oxygen. For the most part they are 92 CHEMISTRY AND DAILY LIFE salts in which glycerine is the base and oleic, palmitic, and stearic acids are the acids. By heating such fats with strong bases like caustic potash or caustic soda glycerine is set free, and the sodium salt of the fatty acid, a soap, is formed. So for example : (C 17 H33COO) 3 C 3 H 5 + 3NaOH = 3 C 17 H 33 COONa glycerine oleate caustic soda sodium oleate cottonseed oil castile soap* + C 3 H 6 (OH) 3 glycerine Mineral fats and oils are hydrocarbons and consequently can- not be transformed into soaps by boiling with lye. It is thus easy to distinguish between mineral oils and those of plant or animal origin. Petroleum products cannot serve as foods. There is no hy- drocarbon which can be used for food. The lower boiling hydrocarbon oils from petroleum are used as fuels, for heating, and for propelling internal combustion engines. When com- pletely burned they form carbon dioxide and water. Hy- drogen burns more readily than carbon, and the richer in hydrogen a hydrocarbon is, the more volatile it is and the more readily it burns. For these reasons petroleum ether, gasoline, naphtha, and benzine are more suitable than kerosene for running so-called gasoline engines, although kerosene is used at present with a fair degree of success, especially in those engines in which that hydrocarbon is transformed into a very fine mist by mechanical means. This mist when mixed with air and subjected to the action of the electric spark in the engine cylinder explodes fairly readily. It should be remembered that all gases which burn in the air will also form mixtures with air which are explosive when ignited. Upon this fact the running of gas and gasoline engines depends. The fuel is in each case volatilized and mixed with air, and the mix- CARBON AND ITS COMPOUNDS 93 ture is then exploded in the cylinder of the engine by ignition. Acetylene is another hydrocarbon which has come into prominent use, particularly for lighting purposes. It is a gas at ordinary temperatures and is readily prepared by the action of calcium carbide, a stonelike solid, upon water, CaC 2 + 2 H 2 O = C 2 H 2 + Ca(OH) 2 calcium carbide water acetylene calcium hydroxide slaked lime The calcium carbide is obtained by heating lime and coke together in the electric furnace out of contact with the air, thus : CaO + 3C CaC 2 + CO lime coke calcium carbide carbon monoxide gas On account of its high carbon content, acetylene has great illuminating power and is consequently frequently used for lighting purposes, especially on automobiles, in running projection lanterns, etc. From petroleum oils, too, gas of high illuminating power may be obtained by " cracking " the oils, which consists of having them come into contact with red-hot stone surfaces out of contact with the air. These oil gases are now produced in many municipal gas plants. They are also put on the market compressed in steel cylinders. Pintsch gas, which is used for lighting railway cars, is a rich oil gas of this char- acter. Blaugas, named from its inventor Blau, is also a sim- ilar oil gas. Hydrocarbons are wry numerous and they may be regarded as the mother substances from which all other compounds of carbon are derived. The simplest hydrocarbon is methane, or marsh gas, CH 4 . It is called marsh gas because it issues from the marshes on warm summer days as the vegetable matter which is always present on the bottom of such waters 94 CHEMISTRY AND DAILY LIFE slowly decays. The gas may also be prepared artificially, by heating sodium acetate with a caustic alkali, thus : CH 3 COONa + NaOH = Na 2 CO 3 + CH 4 sodium acetate sodium hydroxide sodium carbonate methane Marsh gas will burn in the air, thus : CH 4 + 20 2 = C0 2 marsh gas oxygen carbon dioxide Mixed with air it forms an explosive gas mixture, which could be used for running an internal combustion engine. 2H 2 water FIG. 27. Collecting marsh gas from the bottom of a ditch. From marsh gas a variety of products may be obtained. Thus by treatment with chlorine there may be obtained CARBON AND ITS COMPOUNDS 95 successively, methyl chloride CH 3 C1, methylene chloride CH 2 C1 2 , chloroform CHC1 3 , and carbon tetrachloride CC1 4 , viz * : CH 4 - hd 2 = HC1- f- CH 3 C1 CHgCl - hd 2 = HC1 - f- CH 2 C1 2 CH 2 C1 2 H -C1 2 = HC1- h CHC1 3 CHC1 3 H - C1 2 = HCH hCC! 4 Of these compounds chloroform and carbon tetrachloride are by far the most important ones. Both are liquids at ordi- nary temperatures. Chloroform boils at 61 C., is heavier than water, has a rather agreeable odor, and is used as an anaes- thetic. The analogous iodine compound, iodoform, CHI 3 , is a yellow crystalline solid of rather disagreeable odor. It is used in surgery as a dressing for wounds, especially when pus has formed. Carbon tetrachloride is also a colorless, heavy liquid which does not dissolve in water. It boils at 76 and it is not inflammable, which, of course, is to be expected, consider- ing that it consists of carbon and chlorine. Carbon tetra- chloride is an excellent solvent for fats, and hence is often used for removing grease spots from clothes, for which purpose it is to be recommended, for.it does not form dangerous explosive mixtures with the air like gasoline, for example, which is also used in cleaning clothes. However, carbon tetrachloride is much more expensive than gasoline. The trade name "carbon- eum " is sometimes given to carbon tetrachloride. The latter compound is also used in fire extinguishers of various forms , par- ticularly those of the syringe type, which are often sold at exor- bitant prices. Carbon tetrachloride extinguishes fires because it evaporates and crowds the air away from the burning sub- stances, thus making it impossible for combustion to proceed further. During the evaporation of the carbon tetrachloride heat is absorbed, and so the temperature of the burning substances is lowered and this also tends to check the fire. 96 CHEMISTRY AND DAILY LIFE Fires may be extinguished in the following ways : (I) by cutting off the supply of air from the burning materials, (2) by lowering the temperature of the materials on fire so that further combustion is impossible. Now when water is thrown upon an ordinary fire the latter is extinguished for two reasons; namely, (1) the water lowers the temperature, and (2) it also forms clouds of water vapor which crowd the air way. It is obvious that when oils,, which will float on w.ater, are burning, the addition of water cannot prevent access of air to such a fire, for the oils will come on top of the water and continue to burn, the water serving rather as a means of spreading the burning oil. Many incipient fires can be put out by covering them with blankets, rugs, earth, etc., and thus shutting off the air. Burn- ing oils are extinguished by means of carbon dioxide and carbon tetrachloride. Still other chemicals have been suggested for this purpose, but they have not come into common use. Carbon dioxide is one of the very best means known for putting out small fires. It may be delivered from cylinders containing com- pressed carbon dioxide, or may be gener- ated by the action of sulphuric acid upon a saturated solution of baking soda, sodium bicarbonate, NaHCO 3 . The fire extinguish- ers that are to be inverted and are then ready for use, Fig. 28, contain a solution of sodium bicarbonate and a glass bottle in which sulphuric acid has been placed. This bottle is loosely stoppered with a lead stopper. On inverting the entire appa- ratus the sulphuric acid bottle is inverted, its stopper drops out, the acid mingles with the sodium bicarbonate solution, carbon dioxide is generated, and the pressure that thus results FIG. 28. A fire ex- tinguisher partly opened to show its construction. CARBON AND ITS COMPOUNDS 97 is sufficient to cause a stream of the saturated effervescing solution to be delivered from the nozzle. This directed upon the base of the flames cools the burning material and at the same time the carbon dioxide liberated crowds the air away and thus extinguishes the fire. It need hardly be stated that small fires should be promptly put out, and the means for doing so should be on hand wherever there is special danger. Large conflagrations are at times well-nigh impos- sible to cope with because of the mass of water that would be required to prevent access of air and to lower the high FIG. 29. A fire extinguisher in action. temperatures below the kindling point. In such cases the best that can be done is to prevent the spread of the fire by removal of combustible material to which there is imme- diate danger that the flames will spread. When methyl chloride, CH 3 C1, is treated with caustic soda, common salt splits off and simultaneously methyl hydroxide, also called methyl alcohol or spirits of wood, is formed, thus : CH 3 C1 + NaOH = NaCl + CH 3 OH methyl chloride caustic soda common salt wood alcohol H 98 CHEMISTRY AND DAILY LIFE Methyl alcohol is called wood alcohol because it is commer- cially produced by the dry distillation of wood, that is to say by heating wood out of contact with the air, in which process a variety of other products like pyroligneous acetic acid, creosote, etc., are also formed. Methyl alcohol is poisonous. It is used for fuel and for making various chemical substances. So, for instance, upon oxidation of methyl alcohol formalde- hyde is obtained, thus : CH 3 OH + O = HCOH + H 2 methyl oxygen formaldehyde water alcohol Formaldehyde is a pungent, poisonous gas which is soluble in water. Its If) per cent solutions are sold on the market as formaline. It serves as an antiseptic in fumigating rooms, and its solutions serve in treating seeds to free them from injurious fungi, etc., before planting. So, for example, po- tatoes that are infested with potato scab are advantageously treated with formaldehyde solution before planting. The use of formaldehyde as a preservative in milk and other foods and drinks is forbidden by law because the substance is poisonous. On further oxidation of formic aldehyde, formic acid, a fairly powerful acid, may be obtained, thus : HCOH + O = HCOOH formaldehyde oxygen formic acid But one of the hydrogens of this acid is replacable by metals, and when this has been accomplished the resulting salts are the formates. Formic acid was formerly made by the dis- tillation of red ants, in which it occurs. Ordinary alcohol, also called grain alcohol or spirits of wine, is made commercially by the fermentation of sugar by means CARBON AND ITS COMPOUNDS 99 of yeast. As the yeast plant, Fig. 30, lives it produces carbon dioxide from the sugar, thus : grape sugar yeast = 2 CO 2 carbon dioxide 2 C 2 H 5 OH alcohol FIG. 30. Yeast cells. Like wood alcohol, grain alcohol may serve as a fuel. Both substances burn with a hot, blue flame, forming carbon dioxide and water. On careful oxidation ordinary alcohol forms first acetic aldehyde and then acetic acid, thus : CH 3 . CH 2 . OH ordinary alcohol + o = CHaCOH -f acetic aldehyde H 2 O water CHaCOH acetic aldehyde + o = oxygen CHaCOOH acetic acid The oxidation of alcohol to acetic acid is accomplished com- mercially by means of an organism called mother of vinegar. The dilute alcohol is allowed to trickle over beech wood shavings contained in a vat, and in the presence of air and the acetic acid organism the alcohol is oxidized to vinegar. The latter is a solution of about 4 per cent acetic acid. From 100 CHEMISTRY AND DAILY LIFE this by the action of bases a series of salts called acetates is formed. Of the metallic acetates lead acetate, Pb(CH 3 CO 2 )2, is perhaps the most common in practice, though sodium ace- tate, Na . CH 3 CO 2 , and copper acetate, Cu(CH 3 CO 2 )2, are also FIG. 31. Making vinegar by the quick vinegar process. often used. The latter, it will be recalled, forms Paris green when united with copper arsenite. Vinegar is also formed by the fermentation of sugar and subsequent oxidation of the result- ing alcohol in fruit juices, like cider, for example. By adding to cider an antiseptic like boric acid, benzoic acid, or sodium benzoate, fermentation may be prevented, but as these substances are deleterious to health, the practice is not to be recommended. Cider or grape juice may be kept unfermented CARBON AND ITS COMFOT^SS \ ^^ if it is carefully heated to destroy the microorganisms it con- tains, and then, while still hot, sealed in air-tight bottles. Denatured alcohol is grain alcohol which has been rendered unfit for drinking purposes by adding about 10 per cent of wood alcohol, pyridine, C 5 H 5 N, or other poisonous liquid. Denatured alcohol is sold duty-free and is used as a fuel and for manufacturing purposes. Whiskeys and rum contain from 45 to 65 per cent of alcohol, wines from 8 to 20 per cent, and beers from 3 to 5 per cent. As beverages, alcoholic liquids are now recognized as neither necessary nor desirable for the best of health. There are many other alcohols besides wood alcohol and grain alcohol; however but few of them are in common use. Glyc- erine is really an alcohol. Its composition is expressed by the formula C 3 H 5 (OH) 3 . All alcohols contain the OH group. They may all act as bases, reacting with acids to form salts and water. So, for instance, C 2 H 5 OH + CH 3 COOH = CH 3 COOC 2 H 5 + H 2 O alcohol acetic acid ethyl acetate water It will be seen at once that this reaction is similar to that when sodium hydroxide acts on acetic acid : NaOH + CH 3 COOH = CH 3 COONa + H 2 O sodium hydroxide acetic acid sodium acetate water Salts, like ethyl acetate, in which there is a hydrocarbon radical which plays the role of a metal in that it has basic properties, are called ethereal salts or esters. Esters are quite common in nature. So, for example, wintergreen oil, methyl salicylate, is a salt of salicylic acid, in which methyl, CH 3 , the radical from wood alcohol, acts as a base. Banana oil, amyl acetate, CH 3 COO . C 5 Hn, is formed when acetic acid CHEMISTRY AND DAILY LIFE and amyl alcohol react with each other, thus : CH 3 COOH + C 5 H U OH = CH 3 COOC 6 Hii + H 2 O acetic acid amyl alcohol amyl acetate water The fats are salts in which stearic, palmitic, and oleic acids act as acids and glycerine acts as the base. So in mutton tallow we have principally glycerine stearate, in palm oil chiefly glycerine palmitate, while hog's lard, cotton seed and olive oils are rich in glycerine oleate. In general the fats are mixtures of these three salts of glycerine. The softer the fat, the richer it is in oleine or glycerine oleate; the harder the fats, the richer they are in stearine or glycerine stearate. On heating the fats with caustic soda or caustic potash the corresponding salts of sodium or potassium are formed and glycerine is simultaneously produced. These sodium salts of stearic, palmitic, and oleic acids are the hard soaps. The corre- sponding potassium salts are the soft soaps. Glycerine is a product of the process of saponification of fats, for example : (C 17 H 35 COO) 3 C 3 H 5 +3 NaOH=3 Ci 7 Ht 6 COONa-f ej glycerine caustic soda sodium stearate glycerine stearate lye hard soap When used in hard water, soaps form an insoluble curdy mass. This is really the calcium soap. Hard water contains cal- cium salts which react with soap, forming the insoluble calcium soap and a soluble sodium salt, thus : 2 CnHsfiCOONa + CaSO 4 = (Ci7H 35 COO) 2 Ca + Na 2 S0 4 sodium stearate calcium sulphate calcium stearate sodium sulphate When hard water is first treated with washing soda, Na 2 CO 3 , it may be " cleansed " ; that is to say, the calcium may be pre- cipitated in the form of an insoluble compound, calcium carbonate, CaCO 3 , and then the water will not form a calcium soap when treated with ordinary soap. So, for example : CaSO 4 + Na 2 CO 3 = Na 2 SO 4 + 'CaC0 3 calcium sulphate sodium carbonate sodium sulphate calcium carbonate CARBON AND ITS COMPOUNDS 103 Instead of washing soda, borax may also be employed to soften the water, in which case calcium borate is precipitated. Am- monium carbonate may also serve for the same purpose. It is best, however, to use rain water or distilled water for washing purposes and so avoid the use of soda, etc., for the process of cleansing water is always expensive, and the salts that remain in the water are hard on the clothes and also on the hands of the laundry workers. Ether is made by very carefully conducting alcohol into concentrated sulphuric acid heated to about 140 C. Under these conditions alcohol loses water, which is taken up by the sulphuric acid, thus : 2C 2 H 5 OH + H 2 S0 4 = H 2 S0 4 .H 2 O + (C 2 H 5 ) 2 O alcohol sulphuric acid sulphuric acid ether hydrate Ether is the oxide of ethyl, whereas alcohol, it will be recalled, is the hydroxide of ethyl. Ether boils at 35 C. and has a specific gravity of 0.736. It is very inflammable. In this respect it is even more dangerous than gasoline. Ether is used in medicine as an ancesthetic. It is also employed as a solvent for fats and oils, especially in analytical chemistry. Sulphur ethers contain sulphur where ordinary ethers contain oxygen. Sulphur ethers have extremely nauseating odors. Carbolic acid or phenol, C 6 H 5 OH, is really not an organic acid at all. It is an hydroxide derived from benzene, C 6 H 6 , a hydrocarbon present in the light oil obtained by distilling coal tar produced in the gas works. In reality, then, carbolic acid is closely related to alcohols, but since its hydroxyl hydrogen is easily replaced by metals forming so-called phenolates, phenol is commonly called carbolic acid. It is a solid whose crystals melt at 42 C. It turns pink on exposure to air. It is poisonous and corrosive to the skin. Fifteen parts of cold water dissolve about 1 part of phenol, and the solu- 104 CHEMISTRY AND DAILY LIFE tion serves as an antiseptic. Phenol has a very characteristic odor. Creosote is a mixture of guajacol, C 6 H 4 (OCH 3 )OH, and creosol, C 6 H3(CH3)(OCH 3 )OH, produced when wood is heated with but limited access of air, as in smoking fish, hams, bacon, etc. These phenols penetrate into the meats and so preserve them from the ravages of microorganisms, thus preventing putrefaction. Hydrochinone, C 6 H 4 (OH) 2 , and pyrogallol, C 6 H 3 (OH) 3 , also called pyrogallic acid, are also phenols. They are used as developers in photography. From the hydrocarbon benzene, C 6 H 6 , are derived many other useful substances, such as aniline C 6 H 5 NH 2 , nitroben- zene or oil of mirbane, C 6 H 5 NO 2 , benzoic acid C 6 H 5 COOH, whose sodium salt, sodium benzoate C 6 H 5 COONa, has been used as a preservative for foods. Practically all the dye- stuffs that are now used for dyeing fabrics of all kinds are derivatives of benzene, CeHe. They are consequently often spoken of as coal tar dyes, aniline dyes, or coal tar coloring matters. They are excellent, and have practically completely replaced vegetable coloring matters in the arts. Among the many other organic acids that exist there are the following, which are frequently met in daily life : (1) oxalic acid (COOH) 2 , which is made from sawdust by oxi- dizing the latter with the aid of nitric acid. It serves for re- moving ink and rust spots from floors and fabrics, and is fre- quently used as a reducing agent in the laboratory. (2) Lac- tic acid, which occurs in sour milk. It has the formula C 2 H 4 (OH) . COOH. With bases it forms lactates. It is used with baking soda in making biscuits, pancakes, etc. Its silver salt is used as an antiseptic in medicine. (3) Malic acid, CH 2 COOH . CH . OH . COOH occurs in sour apples, in mountain ash berries, and in many other fruits. (4) Tar- taric acid (CH . OH . COOH) 2 occurs in grapes. Its acid potassium salt (CH . OH . COOH)(CH . OH . COOK) is cream CARBON AND ITS COMPOUNDS 105 of tartar and is used in baking powders. (5) Citric acid (CH 2 COOH) 2 . CH . OHCOOH occurs in lemons and other citrus fruits. (6) Butyric acid, C 3 H 7 COOH, is contained in rancid butter and gives to the latter its disagreeable odor. (7) Valeric acid, C 4 H 9 COOH, is contained in the catnip plant, and it is the odor of this acid that is so much esteemed by cats. (8) Hippuric acid, C 6 H 5 . CO . NH . CH 2 . COOH occurs in the urine of herbivorous animals. The so-called carbohydrates form a large and extremely important group of compounds of carbon. They all contain only the elements carbon, hydrogen, and oxygen, and more- over the hydrogen and oxygen are always present in the same proportions by weight as in water, whence the name carbo- hydrate. The carbohydrates consist of (1) the celluloses, (2) the starches, (3) the gums and dextrines, (4) the sugars. Wood, cotton, straw, hemp, linen, etc., when burnt leave behind a certain amount of ash, but aside from this, which represents their mineral content, they consist of almost pure cellulose. Cellulose, then, is one of the most widely distributed compounds in nature, being the material out of which the cell walls of all plants are made. The composition of cel- lulose is expressed by the formula (CeHioOs)^ Cellulose is insoluble in water, and when completely burned forms carbon dioxide and water. In the form of hay, cornstalks, straw, etc., cellulose serves as a food for cattle, horses, and other herbivores, but human beings and carnivorous animals can not use this material for food. It is evident, however, that for fuel, clothing, and shelter, cellulose in various forms is invaluable to mankind. Paper is made from rags, straw, wood pulp, etc. ; it therefore consists of fibers of cellulose that are matted together. Filter paper and blotting paper are unsized papers ; whereas writing paper and many other papers that have a hard, smooth, finished surface contain rosin which has been 106 CHEMISTRY AND DAILY LIFE melted and rolled into the sheets by means of hot rollers. The surface thus finished does not absorb ink so readily, and hence serves well for writing purposes. Nitrates of cellulose may be formed by treating cellulose with a mixture of nitric and sulphuric acids. These nitro- celluloses are true esters of cellulose, for on saponifying them with caustic soda, sodium nitrate and cellulose are formed. Guncotton is cellulose hexanitrate, C^Hi^NOs^O^ It looks much like ordinary cotton, but it is harsher to the touch and on ignition it burns very rapidly and quietly, forming no smoke. Guncotton can be made to explode violently by means of a cap of fulminating mercury. The latter is a dangerous white powder that results when mercury is acted upon by nitric acid in presence of alcohol. Guncotton, then, serves as an explosive. It may be used alone or together with nitro- glycerine. The latter is a viscous yellowish liquid. It is the nitrate of glycerine and is produced when glycerine is acted upon by a mixture of concentrated nitric and sulphuric acids. When absorbed in diatomic earth and molded into sticks, it forms dynamite. Nitroglycerine, like guncotton, may be exploded by means of fulminating mercury, but it may also explode from other shocks and it is consequently a very dangerous compound. With the aid of vaseline and a solvent like acetone, (CH 3 ) 2 . CO, guncotton and nitro- glycerine are worked into threads which are used as smokeless gunpowder. Guncotton, nitroglycerine, and dynamite are used in the arts for blasting purposes. They have also sup- planted black gunpowder in military operations. While gun cotton is not soluble in a mixture of alcohol and ether, cellulose tetranitrate and pentanitrate are soluble in this solvent and form a viscous solution that is known as collodion. It is used for making films in photography, and in medicine it is employed in dressing wounds. On evaporation of the CARBON AND ITS COMPOUNDS 107 alcohol and ether the nitrocellulose remains as a tough transparent film. Celluloid is made of guncotton and camphor. It is therefore a highly combustible material, and this fact should always be borne in mind in using it, so that serious accidents may not result on subjecting it to heat and espe- cially to flames. Starch, too, consists of carbon, hydrogen, and oxygen, and these are present in the same proportion by weight as in cellulose ; but they are combined in a different way than in the latter. Starch, then, has the formula (CsH.i O b ) x . It will not dissolve in water, but when a little of it is rubbed fine in about twice its volume of cold water and then ten to twenty times that volume of boiling water is added to the mix- ture with constant stirring, starch paste results which is used for starching clothes ; it also frequently serves as an adhesive. When starched clothes are afterwards ironed, a gloss is pro- duced because under the hot iron the starch is converted into dextrine which gives the luster to the linen fibers. Starch occurs in all plants, especially in potatoes and other tubers and fleshy roots, but also in grains like rice, wheat, rye, oats, barley, corn, etc. The supply of starch which all seeds contain serves to nourish the young plant till its roots and tops have developed sufficiently to draw sustenance from the soil and the air. In this process starch is gradually changed to sugar, and the latter, being soluble in water, is utilized by the grow- ing embryonic plant. Starch is a most important article of food for man and animals as well. Flour consists of about 70 per cent starch, together with 10 per cent gluten (which is a nitrogen-bearing substance that is akin to the albumins found in the white of egg and hence is also valuable as food) and small quantities of mineral matter (i.e. ash), water, and sugar. When heated to 210 C. starch is converted to dex- trine, C 6 HioO 5 , which is a colorless amorphous substance. On 108 CHEMISTRY AND DAILY LIFE treatment with water it yields a sticky mass, and hence is very commonly used as a cheap adhesive gum. On heating starch with dilute sulphuric acid, it is converted into glucose. The sulphuric acid may afterwards be removed by treatment with lime, for thus insoluble calcium sulphate is produced. Sugars may be divided into two classes, the monoses and biases. The monoses have a composition corresponding to the formula C 6 Hi 2 O 6 , and the common representatives of this group are (1) glucose or dextrose, also known as grape sugar, (2) levulose, or fruit sugar, also called fructose. The com- position of the bioses corresponds to the formula Ci 2 H 22 On, and the chief common representatives of this class are (1) cane sugar, also called saccharose or sucrose, (2) maltose, or malt sugar, and (3) lactose, or milk sugar. Glucose is found in the juice of many fruits, especially of grapes, whence the name grape sugar. Its solutions turn the plane of polarized light to the right, hence this sugar is also called dextrose. When yeast is added to dilute solutions of dextrose, fermentation takes place and alcohol and carbon dioxide are formed, thus : C 6 H 12 O 6 (+ yeast) = 2 C 2 H 5 OH + 2 CO 2 glucose alcohol carbon dioxide Solutions of glucose precipitate red cuprous oxide, Cu 2 O, from hot Fehling's solution. The latter is a strongly alkaline solution of copper sulphate and Rochelle salt, sodium potas- sium tartrate. This solution is much used in testing for re- ducing sugars. For example, by this means in the case of diabetic patients sugar may be detected in the urine. By heating starch with dilute sulphuric acid, glucose is prepared on a large scale, as already mentioned in connection with the consideration of starch. Thousands of tons of glucose are thus manufactured from corn starch annually in the United CARBON AND ITS COMPOUNDS 109 States. Glucose is a good, wholesome food. It is mainly used in candies, table sirups, etc. It is about three-fifths as sweet as cane sugar. Levulose or fructose is also found in the juice of fruits. It also occurs in honey. Its solutions turn the plane of polarized light to the left, whence the name levulose; they also reduce Fehling's solution yielding a precipitate of cu- prous oxide. Furthermore this sugar, too, may be fermented with yeast, yielding alcohol and carbon dioxide. By treating cane sugar solutions with dilute acids both dex- trose and levulose are produced, thus : n -f- H^O = CeH^Oe + CeH^Oe cane sugar water dextrose levulose This process proceeds quite rapidly on heating the solution gently to boiling. Since the resulting liquid turns the plane of polarized light slightly to the left, whereas the original cane sugar solution turned the plane of polarized light to the right, the process of thus changing cane sugar to dextrose and levulose is commonly termed the inversion of cane sugar. The latter compound does not alter Fehling's solution, whereas it will be recalled that both dextrose and levulose effect its reduction. Cane sugar is also known by the names sucrose and saccharose. Sugar cane contains from 15 to 20 per cent of it, whereas sugar beets commonly yield from 10 to 20 per cent. In sorghum, in maple sap, in nuts, and in the blossoms of many plants saccharose abounds. Indeed, it is very widely dis- tributed in the vegetable kingdom. Rock candy is almost pure saccharose and exhibits the beautiful monoclinic prisms in which this substance crystallizes. In water it is very soluble. One part of water will dissolve three times its weight of sugar at room temperature. On carefully heating cane 110 CHEMISTRY AND DAILY LIFE sugar it may be melted to a colorless liquid at 160 C. ; on cooling this yields so-called barley sugar, an amorphous glassy substance, which after a time becomes crystalline. Caramel is a brown substance which is obtained by heating sucrose FIG. 32. A field of sugar beets. to about 200 C. Water is given off as the caramel forms. The latter is much esteemed in candies. Yeast does not cause fermentation of cane sugar solutions at all readily. After a time, however, fermentation with production of alcohol and carbon dioxide does occur, for yeast contains a so-called enzyme, invertase, which slowly inverts sucrose to glucose and levu- lose, and these then ferment. In manufacturing sugar, the juice obtained from sugar cane or sugar beets is treated with about a 1 per cent solution of slaked lime. This con- verts all acids present to calcium salts, serves to coagulate CARBON AND ITS COMPOUNDS 111 proteins present, and at the same time, being an antiseptic, it guards against fermentation. The excess of lime is then removed by means of carbon dioxide, and the solution is decolorized by filtering through bone black, after which it is FIG. 33. A field of sugar cane. evaporated in so-called vacuum pans that are heated by means of steam. On cooling, crystals of sugar separate out. These are whirled in a centrifuge to remove the adhering brown liquid in which they grew. The latter is sold as molasses. Bagasse is the name given to the residue of the beets or cane after the removal of their juice. The bagasse is used as fuel, made into paper, or fed to cows. The annual production of 112 CHEMISTRY AND DAILY LIFE sugar from beets and cane is about ten millions of tons, which shows that it is used very extensively as a food. Maltose or malt sugar occurs in malt. Its solutions turn the plane of polarized light to the right, reduce Fehling's FIG. 34. Vacuum pans used in a sugar factory. solution, and readily ferment with yeast yielding alcohol and carbon dioxide. Maltose is an intermediate product in the formation of alcohol from starch, for when diastase, an enzyme contained in malt, acts upon starch, maltose is produced. This then may be fermented to alcohol and carbon dioxide. Upon these facts the commercial production of fermented liquors like beer depends. The dilute alcoholic solutions obtained by fermentation may be concentrated by fractional distilla- tion, thus yielding whisky, gin, rum, and even alcoholic solutions of 80 per cent and still higher. Alcohol that is entirely free from water cannot be obtained by fractional distillation. When such so-called absolute alcohol is re- CARBON AND ITS COMPOUNDS 113 quired, quicklime is added to alcohol that has been con- centrated by distillation as far as it is profitable to do so by this means. The lime unites with the water present, but not with the alcohol, and the latter can then be distilled off. FIG. 35. A battery of centrifuges used in drying sugar in a sugar mill. When maltose is treated with dilute acids, dextrose only is formed. Lactose or milk sugar occurs in the milk of mammals. Cow's milk contains about 5 per cent of lactose. It is not as sweet as cane sugar and far less soluble in water than the latter. Six parts of water dissolve about 1 part of lactose at room temperature. Lactose turns the plane of polarized light to the right and reduces Fehling's solution, but not as readily as maltose. Toward yeast it acts much like cane sugar, that is to say, fermentation proceeds but very slowly indeed. In this case the products formed, however, are alcohol and lactic acid. Yeast that is quite pure does not ferment lactose at all. 114 CHEMISTRY AND DAILY LIFE The proteins, formerly called proteids, consist of carbon, hy- drogen, oxygen, nitrogen, and sulphur. They are an exceedingly important class of substances, for ivhen the fats, mineral matter, and water are taken from the body of an animal what is left con- sists entirely of proteins. No animal can live without proteins as food. The composition of the proteins is about as follows on the average : carbon, 50.5 to 54.5 per cent ; oxygen, 21 to 24 per cent ; nitrogen, 15 to 17.7 per cent ; hydrogen, 6.5 to 7.3 per cent ; sulphur, 0.3 to 2.3 per cent ; phosphorus, 0.4 to 0.8 per cent. The nucleoproteins so called often con- tain 5 to 6 per cent of phosphorus. The chemistry of the protein bodies is quite complicated, and only of recent years has notable progress toward a better comprehension of the nature of these substances been made. Albumins are pro- teins, they occur, for instance, in eggs, in muscles, in milk, in blood serum, in the seeds of plants, especially in beans, peas, and other legumes. The albumins may be coagulated by means of heat, also by means of acids. The name albumi- noids is given to a class of protein bodies that are closely re- lated to the albumins. So, for example, gelatin, keratin, and elastin are albuminoids. Keratin is the main constituent of hair, hoofs, horns, feathers, nails, cuticle, etc., while elastin is found in the connective tissues. Peptones are formed when albumins are acted upon by pepsin, an enzyme, or so- called unorganized ferment. In the change from albumin to peptone, water is added to the compounds by the action of pepsin. This process takes place in the stomach and is important in the digestion of nitrogenous foods. When proteins are left to putrefy, poisonous substances called ptomaines are produced, among which putrescine C 4 H 8 (NH 2 )2 and cadaverine C 5 Hi (NH 2 )2 may be mentioned. Poisoning caused by eating partially decayed meat, fish, or other animal food is generally due to ptomaines. CARBON AND ITS COMPOUNDS 115 In plants certain non-nutritive nitrogenous constituents occur, which are of importance in common life and ought consequently to be mentioned here. These substances consist of carbon, hydrogen, oxygen, and nitrogen. They have the common property that they are basic in character, that is to say, with acids they form salts. Hence these substances are called alkaloids. Only a few of the most im- portant will be mentioned. Unless otherwise stated they form crystalline solids. Quinine C2oH 24 O 2 N 2 and cinchonine Ci 9 H 2 2ON 2 occur in Peruvian bark. The sulphate of quinine is commonly used in treatment of malaria. Morphine CnHigOaN, codeine Ci 8 H 2 iXO 3 , and narco- tine C22H 2 3O 7 N occur in opium, the dried sap of the partially ripe pods of the opium poppy. Mor- phine is used to produce sleep and is commonly prescribed as the hy- drochloride. Strychnine C 2 iH 22 O 2 N 2 and brucine C 2 3H 26 N 2 O 4 are found in mix vomica. These are extremely bitter and wry poisonous, producing death with concomitant convulsions and final muscular rigor. Nicotine CioHi 4 N 2 is found in tobacco. It is a poisonous liquid which boils at 247 C. Cocaine CnHiC^N occurs in coca leaves and is used as a local anaesthetic. Atropine Ci7H 23 O 3 N is found in nightshade, the deadly Atropa belladonna. Oculists use this alkaloid to produce expansion of the pupil of the eye. FIG. 36. An oriental poppy. From this plant opium is obtained. 116 CHEMISTRY AND DAILY LIFE QUESTIONS 1. Make a list of the various forms in which carbon occurs. How do we know that all of these are the same element carbon? 2. Mention the uses of charcoal, also of bone black. 3. What are the different varieties of coal? What is the essential difference between them ? 4. How is coal gas made ? What does it contain ? 6. Explain how it is that ammonia is a product of the gas works. 6. How is a flame produced ? Give an example. 7. What causes some flames to be luminous and others to be non- luminous ? Give examples of luminous and non-luminous flames. 8. Draw a diagram of a Bunsen burner and explain it. 9. What are the properties of carbon dioxide? How is it formed ? How does it occur in nature ? What is it used for ? 10. How is carbon monoxide formed ? What are its chief char- acteristics ? 11. What is water gas ? How is it made ? 12. How is starch formed ? How many pounds of carbon dioxide would have to be decomposed to form 5 pounds of starch ? 13. What is the so-called carbon cycle ? 14. What is soda water ? 15. What is lime, and how is it produced ? 16. Explain the process of slaking lime. 17. What happens when lime mortar sets ? 18. What is petroleum ? Mention the various important petro- leum products that are on the market. 19. How can a mineral oil be distinguished from one that is of plant or animal origin ? 20. What is acetylene ? How many grams of calcium carbide are necessary to make 100 liters of acetylene at and 760 mm. pressure ? 21. What is the difference between a hydrocarbon and a carbo- hydrate ? Name five examples of each. 22. What is chloroform ? For what purpose is it used ? 23. What is carbon tetrachloride ? What are its uses ? 24. How may fires be extinguished ? 25. Why should water not be used to extinguish burning oils ? CARBON AND ITS COMPOUNDS 117 26. What is wood alcohol ? What are its uses ? 27. How much wood alcohol would be required to produce two pounds of pure formaldehyde ? 28. What is formaline, and what is it used for ? 29. How is ordinary alcohol produced ? * What are its uses ? 30. How much grape sugar would be required to produce 100 pounds of pure alcohol ? 31. What is vinegar ? How is it produced ? 32. What is denatured alcohol ? Why is it made ? 33. What is glycerine ? How is it produced ? 34. How are alcohol and ether related chemically ? 36. To what class of substances do the plant and animal fats and oils belong ? 36. What is hard soap ? Soft soap ? 37. What is meant by the term "hard water"? How is water softened ? 38. What is carbolic acid ? What is it used for ? 39. Why does it preserve hams and fish to smoke them ? 40. What is benzene ? Benzoic acid ? Aniline ? 41. Mention six organic acids .and tell where they occur. 42. What substances are called carbohydrates and why ? Where do these occur in nature ? 43. Which of the carbohydrates serve as human food ? Classify the carbohydrates. 44. To what animals may cellulose serve as food ? Will cellulose in any form do for this purpose ? Explain by a few examples. 45. What is guncotton? Dynamite? Smokeless powder? Celluloid? 46. What are the properties of starch ? How test for the pres- ence of starch ? 47. What is flour? 48. Classify the sugars, and give the chief characteristics of each of the principal sugars. 49. What is a protein ? Give several examples. 60. To what class of bodies do morphine, quinine, and strychnine belong ? CHAPTER X THE METALS OF THE ALKALIES AND THE ALKALINE EARTHS. THE metals of the alkalies are potassium, sodium, lithium, ccesium, and rubidium. The last two are of such rare occur- i rence that they will not be con- sidered here. Even lithium is found only in small amounts. It occurs in certain micas and some- times in a few plants like tobacco and beets. In general its proper- ties, and also those of caesium and rubidium, are similar to the prop- erties of potassium and sodium. The metals of the alkalies never occur in nature in the free or un- combined state. They are only found with other elements in the form of salts. Plant ashes always contain po- tassium carbonate, K^COa, also called potash. It has long been the practice to leach out wood ashes with water, and use the potash solution thus obtained for the purpose of making soap. Potassium is an element that is essential to plant and animal life. Without potassium salts plants cannot grow. All soils 118 FIG. 37. Buckwheat. On the right all the elements are pres- ent ; on the left all except po- tassium. THE METALS OF THE ALKALIES 119 contain potassium in the form of soluble salts, although in most cases in insufficient quantity. For this reason, it has long been the practice to return potassium salts to the land. Certain quantities get back through the application of stable manure. Wood ashes are commonly applied to fields when pot- ash is lacking. Feldspar, which is a silicate of potassium and FIG. 38. Geological formation of the Stassfurt salt beds. aluminium, occurs in all granitic rocks and in many clay soils, and plants secure notable amounts of potash from this source, for the soil waters gradually extract potash from the feldspar, though this process is quite slow. The largest bed of potassium salts found in nature is that at Stassfurt in Germany. Here potassium occurs chiefly as carnallite KC1 . MgCl2 . 6 H 2 O and as kainite KC1 . MgSO 4 . 3 H 2 O. It is also found there as sylvite, muriate of potassium, KC1, to some extent. The potassium salt deposits at Stassfurt form layers varying in thickness from sixty to one hundred feet. They rest upon thick beds of common salt. Practically the entire supply of potassium salts of the world is derived from this enormous 120 CHEMISTRY AND DAILY LIFE deposit. However, in small amounts potassium salts are found very widely distributed. So potassium occurs not only in all plants, but in bones, muscles, blood, milk, albumin, and the various secretions of animals. The earth's crust contains about 2.45 per cent potassium, and in oceanic water there is about 0.04 per cent. Caustic potash, potassium hydroxide, KOH, may be obtained by treating potassium carbonate with slaked lime, thus: K 2 CO 3 + Ca(OH) 2 = CaCO 3 + 2 KOH potassium calcium calcium carbonate caustic potash or carbonate hydroxide (insoluble) potassium hydroxide Large quantities of potassium hydroxide are now prepared by electrolysis of potassium chloride, muriate of potash. Thus this salt is decomposed into chlorine and potassium, and when the latter acts on water, hydrogen and potassium hydroxide are formed ; the reactions involved are therefore : 2 KC1 by electrolysis = 2 K + C1 2 and 2 K + 2 H 2 =2 KQH + H 2 By boiling down caustic potash solutions (which must be done in silver or iron vessels, for glass or porcelain would be strongly attacked by the lye) solid potassium hydroxide may be obtained. It is a white, hard, brittle solid which dissolves very readily in water with evolution of considerable heat. Its solutions corrode and disintegrate animal and vegetable tissues, hence the name caustic potash. It is used in making soft soap, being placed on the market as potash lye. On neu- tralization with acids, it forms various potassium salts. A few of the most important ones will now be described. Potassium carbonate has already been mentioned. While it was formerly derived almost entirely from wood ashes, it is now made from the Stassfurt salts by a method similar to THE METALS OF THE ALKALIES 121 that of preparing sodium carbonate by the so-called LeBlanc process (which see). Potassium carbonate is also known commercially as pearlash. It is employed in the manufacture of soft soap, hard glass, and various salts of potassium. Saltpeter, potassium nitrate, KNO 3 , is present in small amounts in all soils. It is a white crystalline salt. At ordinary temperatures 100 parts of water dissolve 13 parts of potassium nitrate, but the same amount of boiling water dissolves 247 parts of this salt. Saltpeter is prepared com- mercially by adding potassium chloride (from the Stassfurt deposits) to hot saturated solutions of Chili saltpeter, NaNO 3 , thus: NaNO 3 + KC1 = NaCl + KNO 3 sodium nitrate potassium sodium chloride potassium chloride nitrate The common salt, NaCl, is not nearly as soluble as the potas- sium nitrate, consequently the latter remains dissolved while the former is precipitated. From the clear solution, crystals of potassium nitrate are then obtained on cooling. Salt- peter is used in the manufacture of black gunpowder, which consists of 75 parts saltpeter, 11 parts sulphur, and 17. 5 parts charcoal. When it is exploded a portion of the latter ingre- dient always remains unoxidized and so forms black smoke. The use of black gunpowder in war is now a thing of the past. Saltpeter is also used in salting meats; so, for example, 60 parts of common salt, 1 part of saltpeter, and 10 parts of sugar, all dissolved in 400 parts of water, make a good solu- tion for salting beef or pork. Potassium chloride, KC1, is a white crystalline salt. It is readily soluble in water and tastes salty. It is chiefly obtained from the Stassfurt deposits and is the commonest and cheapest of the potassium salts. In crude form it serves as a fertilizer. From it caustic potash, saltpeter, and other 122 CHEMISTRY AND DAILY LIFE potassium salts are prepared. This salt is also found in sea water. In the blood and urine of herbivorous animals it is also present in minute amounts. Potassium bromide, KBr, and potassium iodide, KI, crystallize in cubes, which are extremely soluble in water. Both of these salts are used in medicine and in photography. Potassium chlorate, KC1O 3 , is produced by conducting chlorine into hot solutions of caustic potash, thus : 6 KOH + 3 C1 2 = KC10 3 + 5 KC1 + 3 H 2 O caustic potash chlorine potassium potassium water chlorate chloride By electrolyzing hot potassium chloride solutions both caustic potash and chlorine are simultaneously formed, and so most of the potassium chlorate in the market is prepared by this means. Potassium chlorate is used in making oxygen, matches, fireworks, and explosives. It is also used in medicine, particu- larly as a gargle for sore throat. Potassium silicate, K 2 SiO 3 , is made by fusing sand with potassium carbonate, thus : Si0 2 + K 2 C0 3 = K 2 Si0 8 + C0 2 silica potassium potassium carbon dioxide carbonate silicate This salt is soluble in water and its sirupy solutions are popularly called potassium water glass. Potassium sulphate, K 2 SO 4 , is found at Stassfurt in kainite. It is also a constituent of potassium alum. In crude form it serves as a fertilizer. It is employed in the manufacture of hard glass, alum, and potassium carbonate. Potassium phosphate, K 2 HPO 4 , is present in minute amounts in the blood and urine of all carnivores. It is a white crystalline salt which dissolves readily in water. Potassium cyanide, KCN, is an extremely poisonous salt. It THE METALS OF THE ALKALIES 123 is used in photography, in extracting gold from its ores, in gold and silver plating, and also in preparing hydro- cyanic acid gas, HCN, for fumigating trees. Potassium cyanide and hydrogen cyanide (also known as Prussic acid) are terrible poisons and must be used with extreme care. To the Bunsen flame potassium salts impart a beautiful violet color. Metallic potassium itself was prepared in 1807 by elec- trolysis of molten caustic potash by Sir Humphry Davy. Its specific gravity is 0.865 at 15 C. The metal is rather soft, melts at 62.5 C., and boils at 667 C. With water it reacts violently, forming hydrogen and caustic potash, hence the metal must be kept under petroleum oils to protect it from the moisture of the air. Metallic sodium and metallic lithium have properties that are in general similar to those of metallic potassium. They may be obtained by electrolysis of molten caustic soda, NaOH, and caustic lithia, LiOH, respectively. Sodium and lithium are therefore also kept under petroleum tb protect them from the moisture of the air. Sodium is silver-white, somewhat harder than potassium, melts at 95.5 C. and boils at 742 C. At 15 C. its specific gravity is 0.974. All sodium salts are soluble in water. To the Bunsen flame they all impart a yellow color, whereas lithium salts color the flame a characteristic red. In general the sodium salts are similar to those of potassium. The chief source of all sodium compounds is common salt, sodium chloride, NaCl, which occurs in the sea, in the waters of salt lakes and wells, and also in layers in solid form, par- ticularly at Stassfurt and Reichenhall in Germany, in Galicia, in England, in the states of New York, Michigan, Texas, Utah, and California, also in Africa and Asia. Salt is very widely distributed on the globe. It is necessary to animal life, 124 CHEMISTRY AND DAILY LIFE SIR HUMPHRY DAVY, 1778-1829, the discoverer of the alkali metals. THE METALS OF THE ALKALIES 125 and occurs in the blood, the tissues, and all of the secretions. Just as land plants take up potassium chloride, so sea plants take up common salt. Common salt crystallizes in cubes and melts at about 800 C. It is about as soluble in cold as in hot water. At room temperature 100 parts of water dis- solve 36 parts of salt, while the same amount of boiling water dissolves but 39 parts. The United States alone produces nearly four and one-half million tons of salt annually and this is only about one-fourth of the world's yearly production. Salt is used for making chlorine, hydrochloric acid, caustic soda, sodium carbonate, and other sodium compounds. Caustic soda, sodium hydroxide, NaOH, is quite similar to potasium hydroxide in its properties and is prepared in an analogous manner. It is used in making hard soap, also in " softening " water. It is further employed in the manu- facture of carbolic acid and oxalic acid, also in making paper and other articles. Being cheaper than caustic potash, it is used in place of the latter whenever it is possible. Sodium carbonate, Na 2 CO 3 , is also often called sal soda, or soda. Just as potassium carbonate or potash occurs in the ashes of land plants, so soda is found in the ashes of marine plants. Soda is used to make common glass, hard soap, and many other sodium compounds. It is therefore manufactured on a very large scale. There are two processes of manufactur- ing soda, namely, the LeBlanc process and the Solvay process. Both utilize common salt as the raw material. In the LeBlanc soda process salt is first treated with sul- phuric acid, thus : 2 NaCl + H 2 SO 4 = NaCl + NaHSO 4 + HC1 salt sulphuric acid salt cake hydrochloric acid The so-called " salt cake " consists of salt and acid so- dium sulphate. This is then heated in a suitable fur- 126 CHEMISTRY AND DAILY LIFE nace, thus forming more hydrochloric acid, and sodium sulphate: 4 = NaaSO4 + HC1 salt cake sodium hydrochloric sulphate acid The sodium sulphate is then mixed with charcoal and cal- cium carbonate and heated in a rotary furnace ; in this way carbon monoxide escapes and calcium sulphide and soda are formed, thus : Na 2 SO 4 + 4C Na*S + 4 CO sodium sulphate carbon sodium sulphide carbon monoxide and Na 2 S + CaCO 3 = Na 2 CO 3 + CaS sodium calcium soda calcium sulphide carbonate sulphide The entire mass is then treated with water. The sodium carbonate dissolves and calcium sulphide remains as an in- soluble residue. From the clear solution, crystals are ob- tained which have the composition Na 2 CO 3 . 10 H 2 O. This is so-called washing soda or crystallized soda. On heating the crystals they lose water and form calcined soda or an- hydrous sodium carbonate, Na 2 CO 3 . By the Solvay process, sodium bicarbonate, NaIICO 3 , is first obtained. This salt is relatively less soluble than the corresponding ammonium salt, ammonium bicarbonate, NH 4 HCO 3 , which may be obtained when carbon dioxide is passed into strong ammonia water, thus : NH 4 OH + CO 2 = NH 4 HCO 3 ammonium carbon ammonium hydroxide dioxide bicarbonate On treatment of a saturated solution of sodium chloride with ammonium bicarbonate, sodium bicarbonate is precip- itated, thus : NaCl + NH 4 HCO 3 = NaHCO 3 + NH 4 C1 - salt ammonium sodium ammonium bicarbonate bicarbonate chloride THE METALS OF THE ALKALIES 127 The ammonium chloride remains in solution. In actual practice, then, the Solvay process consists in conducting both ammonia and carbon dioxide into a saturated solution of com- mon salt. Thus both of the changes expressed by the last two equations take place together. Sodium bicarbonate is also called saleratus or baking soda. It has a mild alkaline mz M?t ^>P ; ^- Y .v? FIG. 39. A Chilian saltpeter bed. reaction, and is used in medicine, in baking powders, etc. On heating sodium bicarbonate it loses water and carbon dioxide and forms sodium carbonate, thus : 2 XaHCO 3 = Na 2 CO 3 + H 2 O + C0 2 Even when solutions of sodium bicarbonate are heated, car- bon dioxide is evolved ; upon this fact the use of saleratus in baking depends. Sodium nitrate, XuXO 3 , is found in large beds in Chili and is consequently known as Chili saltpeter. With it occur 128 CHEMISTRY AND DAILY LIFE smaller amounts of sodium chloride, sodium sulphate, and sodium iodate, NaIO 3 . The latter compound is the chief commercial source of iodine. Sodium nitrate is used as a fertilizer on account of its nitrogen content. It also is em- ployed in large quantities in making nitric acid, thus : NaN0 3 + H 2 SO 4 = NaHS0 4 + HNO 3 Sodium nitrate should be added to the soil only in small amounts at a time. Sodium sulphate, Na 2 SO 4 , is formed as one of 'the products of the Leblanc soda process. It dissolves readily in water, and when it crystallizes from the latter at room temperatures beautiful crystals of the composition Na 2 SO 4 . 10 H 2 O are formed. This is the so-called Glauber's salt, which is so much used as a laxative. It was first made by Johann Rudolf Glauber in 1658. Sodium sulphite, Na 2 SO 3 , is made by passing sulphur dioxide into sodium hydroxide solution. This salt is used in photography. Sodium thiosulphate, Na 2 S 2 O 3 . 5 H 2 O, is *the so-called " hyposulphite of soda" or "hypo," so much used in making the "fixing bath" in photography. It is produced by boil- ing a solution of sodium sulphite with flowers of sulphur. Sodium silicate, Na 2 SiO 3 , or sodium water glass, is produced by fusing silica with sodium carbonate or with a mixture of carbon and Glauber's salt. In laundry soaps it serves as a " filler." It is also employed in cementing together the fibers of mineral wool and asbestos, and in fireproofing wood, fab- rics, etc. Borax has the composition Na 2 B 4 O 7 . 10 H 2 O. Its uses have already been mentioned. The metals of the alkaline earths are barium, strontium, calcium, and magnesium. They are so called because their respective oxides BaO, SrO, CaO, and MgO are alkaline, THE METALS OF THE ALKALIES 129 which fact is apparent when they are treated with moist red litmus paper, for example. These metals may be pre- pared by the electrolysis of their molten chlorides. They all are silver-white in appearance, and when treated with water they yield hydrogen and the hydroxide of the metal em- ployed. The hydroxides are soluble in water, but not copi- ously. The chief source of barium is barium sulphate, BaSO 4 , also called heavy spar or " barytes." It occurs in nature in fairly pure form in compact masses. Its specific gravity is 4.48. Ground to powder, it serves as a white pigment in paints, being called " permanent white." It is not infre- quently used as a cheap adulterant in white lead, a practice that is to be condemned. Barium sulphate is almost com- pletely insoluble in water and also in dilute acids. It pre- cipitates whenever a solution of a sulphate is treated with barium chloride solution or a solution of any other barium salt, so, for instance : CuSO 4 + BaCl 2 BaSO 4 + CuCl 2 copper sulphate barium chloride barium sulphate copper chloride Barium chloride is consequently very commonly used in testing for sulphates, also for determining the amount of sulphur present in a compound, for the latter process commonly con- sists of converting the sulphur into a soluble sulphate and then precipitating barium sulphate by adding an excess of barium chloride solution. From the amount of barium sulphate formed the amount of sulphur may be computed, for the composition of barium sulphate is well known. Barium nitrate, Ba(NO 3 )2, is used for making green Bengal lights. Barium salts are all poisonous. Cattle in our Western states have been reported as seriously poisoned by eating herbs that had taken up barium from the soil. This 130 CHEMISTRY AND DAILY LIFE is, of course, very unusual, for barium is ordinarily not found in soils. Strontium compounds are entirely similar to those of barium, only they are of rarer occurrence. Strontium ni- trate, Sr(NO 3 )2, is used in fireworks and for making red Bengal lights. All strontium salts color the Bunsen flame red, while those of barium color it green, and calcium compounds produce an orange-yellow color. Calcium carbonate, CaC0 3 , is found in large quantities in nature as limestone, marble, and chalk. Iceland spar or calcite is a very pure crystalline form of calcium carbonate. When limestone occurs mixed with clay, it is called marl. Calcium sulphate, CaSO 4 , occurs as anhydrite, gypsum, CaSO 4 . 2 H 2 O, and alabaster. In natural waters calcium bicarbonate and calcium sulphate are quite commonly found. They produce the hard sediments in teakettles, boilers, etc. These sediments generally also contain silica, oxides of iron, alu- minum, magnesium, etc. These have been dissolved from the rocks or soils with which the waters have come in contact. Marble and limestone serve as building stones, also for making lime, glass, and cement, and reducing iron ores. In the form of chalk calcium carbonate is used as whiting, which when ground up with linseed oil forms putty. Lime, CaO, is made by heating limestone or marble in kilns. Thus, carbon dioxide is driven off and the oxide of calcium remains : CaCO 3 on heating = CaO + CO 2 limestone lime carbon dioxide The slaking of lime and its use in making mortar have already been described. Calcium chloride, CaCl 2 , is formed when hydrochloric acid acts on limestone, thus : THE METALS OF THE ALKALIES 131 CaC0 3 limestone 2HC1 = hydrochloric acid CaCl 2 calcium chloride H 2 water C0 2 carbon dioxide Calcium chloride is very soluble in water, and its brines are commonly used as the cold liquid that cir- culates in the pipes in re- frigerators. This brine is kept cold by means of evaporation of ammonia which has been liquefied under pressure. On carefully heating gypsum, CaSO 4 . 2 H 2 O, to about 110 it loses water and becomes CaSO 4 . | H 2 O, which is known as plaster of Paris. When it is af- terward mixed with water, the whole sets or hardens to a solid mass, which depends upon the fact that plaster of Paris and water unite and again form gypsum, thus : CaSO 4 . J H 2 O + 1} H 2 O = CaSO 4 . 2 H 2 O plaster of Paris water gypsum Plaster of Paris is much used in making hard wall plaster, gypsum, calsomine, casts, and bandages in surgery. As all plants and animals contain sulphur and calcium salts, calcium sulphate in the form of gypsum is commonly used as a fertilizer under the name of land plaster. FIG. 40. A modern lime kiln. 132 CHEMISTRY AND DAILY LIFE Calcium nitrate, Ca(NO 3 ) 2 , is very soluble in water. It is formed in composted manure, and also on the white- washed walls of stables. The ammonia present in the stable air is gradually oxidized to nitric acid, which then attacks the calcium carbonate on the walls, forming calcium nitrate and carbon di- oxide. This, of course, destroys the wall coat- ing. The calcium nitrate in the soil is a good plant food, for it contains both calcium and nitrogen in available form. Calcium phosphate, Ca3(PO 4 ) 2 , is insoluble in water but soluble in di- lute acids. It occurs in crystals as apatite, is the chief constituent of bones, and is found in phosphate rock, guano, and barnyard manure. The fact that by treating bones or phosphate rock with sul- phuric acid, superphosphate fertilizer is formed has already been mentioned. Over two and one-half million tons of phos- phate rock are used annually in the United States as fertilizer. The exportation of phosphate rock should be forbidden by law, for the supply of this important material, without which agriculture cannot thrive, is limited. The beds of phosphate FIG. 41. A broken leg in a plaster of Paris cast. THE METALS OF THE ALKALIES 133 rock occur mainly in Florida and South Carolina, but re- cently deposits have also been found in some of the Rocky Mountain states. As calcium silicate, calcium occurs in many rocks, espe- cially in complex silicates like feldspars, garnet, mica, horn- blende, etc. Portland cement is made by heating limestone together with aluminum silicates in a kiln, and then grinding up the clinker to a fine powder. This is usually grayish or grayish FIG. 42. Concrete fence posts. brown in color. When mixed with water it will unite with the latter, forming a hard stonelike mass which in appearance resembles the native stone found at Portland, England, whence the name Portland cement. No chemical formula can be given for Portland cement, but a good Portland cement must contain silica, SiO2, from 20 to 25 per cent ; lime, CaO, from 58 to 65 per cent; alumina, from 5 to 10 per cent. Magnesia, MgO, must not exceed 5 per cent, and the less 134 CHEMISTRY AND DAILY LIFE there is of it, the better. Iron oxide, Fe 2 O3, may be present in notable quantities without doing harm ; it is commonly found in cements to the extent of from 2 to 5 per cent. Other impurities, like small amounts of soda, potash, and sul- phates, are generally also found, commonly less than 1 per cent. At Milwaukee, Louisville, and Rosendale, N.Y., occur natural deposits of limestone that contain almost the right amounts of aluminum silicates to make Portland cement. This material is heated in kilns, and the resulting clinker is then ground up and sold as " natural cement." While quite serviceable for many purposes and cheaper than Port- land cement, it nevertheless is not equal to the latter in wearing qualities. Portland cement sets, even under water, and hence can be used in damp places where ordinary lime mortar will not harden at all. Mixed with sand and crushed stone or gravel and then saturated with water, Portland cement forms so-called concrete. In modern structures, iron or steel rods or networks of wires are often embedded in concrete to strengthen it. This is then known as reenforced concrete. It is much used in erecting bridges, fireproof buildings, etc. Nearly fifteen thousand tons of Portland cement are produced annually in the United States alone. The slags from blast furnaces, which contain silica, alumina, and lime, are now also used in the manufacture of Portland cement, the lacking ingredients being added, of course, before firing the material in the kiln. The cements from blast fur- nace slag are generally richer in iron and browner in color than other cements. They are of very good quality, nevertheless. Glass commonly consists of a mixture of the silicates of calcium and sodium. It is made by melting together silica, SiO2, limestone, CaCOs, and soda, Na 2 CO3. The final product corresponds approximately to 'the composition THE METALS OF THE ALKALIES CaO . 3 SiO 2 -f Na 2 . 3 SiO 2 135 and is known as soft glass or soda lime glass. It is used in making windows, bottles, and ordinary glass dishes. The mixture is melted in pots of fire clay about four feet high and four feet in diameter. Potash lime glass (A) (B) FIG. 43. (A) Blowing a glass bottle. (B) Making window glass. is hard glass. It is also called Bohemian glass. It contains potassium silicate instead of sodium silicate. It is not as readily attacked by water, acids, and alkalies as soda lime glass, and hence is used for making glassware for chemical purposes, like beakers, retorts, etc. Crown glass is also a potash lime glass. Jena glass contains boric anhydride. Blue glass is colored by cobalt oxide ; brown glass by ferric oxide ; green glass by ferrous oxide or chromic oxide ; etc. Black glass contains large amounts of iron and other metallic oxides. Milk glass contains suspended particles of calcium 136 CHEMISTRY AND DAILY LIFE phosphate. The Egyptians and Phoenicians knew how to make glass long before the Christian era. Window glass was first used in the sixteenth century. Metallic magnesium is sold in the market in powder, and in the form of wire or ribbon. It burns with a brilliant white light, forming magnesium oxide, MgO. The metal is used for fireworks, for flash lights in signalling and in photography. FIG. 44. Manufacture of glass. Rolling out plate glass. Five parts of magnesium powder to nine parts of pulverized potassium chlorate makes a good flash-light powder. In nature magnesium is found as magnesite, the carbonate MgCO 3 , but far more commonly as dolomite or magnesium limestone, MgCO 3 . CaCO 3 , which has already received men- tion. Magnesium occurs in many natural silicates very widely distributed. So there are hornblende Mg 2 CaFeSi 4 Oi 2 , asbestos Mg 3 Si 2 O 7 . 2 H 2 O, meerschaum Mg 2 Si 3 O 8 . 4 H 2 O, and many other complex silicates that contain magnesium. All plants and animals contain magnesium compounds, and THE METALS OF THE ALKALIES 137 these finally appear in the ash as phosphates and carbonates. Magnesium ammonium phosphate occurs in guano and is valuable as a fertilizer. Magnesium oxide and magnesium carbonate are used in medicine. The latter compound also serves as a face powder. Talc, a hydrous magnesium silicate, is used for the same pur- pose. Magnesium sulphate serves as a purgative under the name Epsom salt. A II magnesium salts have a bitter taste and act as laxatives. When found in natural waters, the latter are called bitter waters. Some of them, like the waters of the springs at Epsom in England and Hunyad in Hungary, are quite famous. It is to be borne in mind that while the car- bonates and phosphates of barium, strontium, calcium, and magnesium are difficultly soluble in water, and the sulphate of barium is insoluble, the sulphates of strontium and calcium are sparingly soluble, the sulphate of magnesium is quite readily soluble. Four parts of water dissolve one part of mag- nesium sulphate crystals at room temperatures. In their chemical behavior, magnesium compounds are in many ways analogous to those of zinc. Magnesium chloride, which is cheap, as it is a by-product of the potash salt industry at Stassfurt, forms with mag- nesium oxide a basic magnesium chloride. This sets as a hard mass, and is frequently used as a cement in making floors, wainscoting, etc. The oxide obtained by calcining native magnesium carbonate is mixed dry with the appropriate amount of sand, sawdust, or other filler and the desired pigment, and then this powder is treated with magnesium chloride solution of proper strength, as in making a mortar. The material is spread out with trowels, and sets in about 24 to 30 hours to a hard mass of good wearing qualities. In Germany the material is called " Stemholz." In America this cement is being intro- 138 CHEMISTRY AND DAILY LIFE duced, and various fanciful trade names are being ap- plied to it. The salts of the element radium are similar to those of the metals of the alkaline earths. Radium salts glow in the dark. The light affects photographic plates. The intensity of this light increases with the purity of the radium compounds. Heat is also constantly evolved by radium compounds, and the air in their immediate neighborhood has less electrical resistance than ordinary air. This is due to the radiations or so-called emanations that are continually being emitted by radium compounds. In extremely minute amounts, which indeed are almost infinitesimal, radium compounds are very widely distributed in soils and natural waters. In uranium com- pounds, especially in an oxide of uranium known as pitch- blende, radium occurs in relatively larger quantities. The emanations from radium are able to destroy germs. They also stop the germinating power of seeds and kill the tissues of living beings. Experiments with radium compounds must there- fore be conducted with great care. Their value as an agent for curing diseases is still a question to be determined. QUESTIONS 1. Name the alkali metals and state their general characteristics. 2. What is potash and how is caustic potash prepared from it ? 3. Where is potassium found in nature ? 4. Mention five important compounds of potassium. 6. What is black gunpowder chemically ? 6. How is potassium chlorate made, and what is it used for ? 7. What is potassium cyanide? What are its properties and uses? 8. How may sodium compounds be distinguished from potas- sium compounds? 9. Where is common salt found in nature ? What is salt used for ? THE METALS OF THE ALKALIES 139 10. What are the two processes for making soda on a commercial scale ? What uses are made of soda ? 11. What is the difference between washing soda and baking soda? 12. (a) What is Glauber's salt? (6) What is the "hypo" used in photography ? (c) What is sodium water glass, and for what is it used ? 13. Name the metals of the alkaline earths and state why they are so called. 14. What is barium sulphate ? What is it used for ? By what other names is it known ? 16. How much barium is there in a ton of barium sulphate ? 16. What use is made of strontium nitrate ? 17. What are the various forms in which calcium carbonate occurs? How do we know that these are all one and the same chemically.? 18. What is gypsum ? Plaster of Paris ? 19. What happens when plaster of Paris sets ? 20. Explain the action of hydrochloric acid on limestone. 21. Of what value is calcium nitrate in soils ? 22. What is phosphate rock and why is it valuable ? 23. How is Portland cement made ? 24. What is " natural cement" ? Where is it made ? 26. What are the uses of Portland cement ? 26. What is common glass chemically? How does soft glass differ chemically from hard glass ? How is colored glass produced ? When was window glass first made ? 27. For what purposes is metallic magnesium used ? 28. Mention five compounds of magnesium that occur in nature. 29. What is Epsom salt chemically ? 30. Mention some of the characteristics of compounds of radium. CHAPTER XI ALUMINUM, THE HEAVY METALS, AND THEIR IMPORTANT ALLOYS OF all the metals aluminum is by far the most abundant, forming about 8 per cent of the material of which the earth's crust is composed. It is found mainly as a constituent of all the common siliceous rocks, especially in feld- spars, clays, micas, gran- ites, slates, etc. It never occurs in the uncombined state, for it has great affin- ity for oxygen, and in- deed it nearly always is found in compounds with that element. Emery is an impure form of alu- minum oxide, being colored brown by the presence of oxides of iron. Sapphires and rubies are beautifully crystallized aluminum oxide and are highly prized as gems. Bauxite is a hydrated ox- ide of aluminum, and is used as a source for the manufacture of alumi- 140 FIG. 45. An emery wheel. ALUMINUM, HEAVY METALS, ALLOYS 141 num. In Greenland a double fluoride of sodium and alumi- num, A1F 3 . 3 NaF, is found. It forms white masses insoluble in water. This mineral is called cryolite. It may be melted quite readily and the molten mass dissolves aluminum oxide. By passing the electric current through this molten mass metallic aluminum is deposited and oxygen comes off at the other pole. The bath is replenished by continually feeding in more aluminum oxide, also called alumina. The molten mass is contained in a vessel of graphitic carbon which serves FIG. 46. Cooking utensils made of aluminum. as the pole on which the metal is deposited. Carbon sticks dipped into the molten mass form the other pole. Eight volts electrical pressure is sufficient to keep the process going, but the current strength is usually several thousand amperes. Thus the heat developed by the electric current is itself sufficient to keep the bath and the deposited metal as well in the molten 142 CHEMISTRY AND DAILY LIFE state, once the process has been started. The molten alu- minum is tapped off from the bottom of the vessel from time to time. Over twenty-five thousand tons of aluminum are thus produced annually in the United States alone. The metal sells for less than twenty cents per pound. Aluminum metal is silver-white in appearance, though on exposure to the air it oxidizes slowly, and the film of oxide on its surface gives the metal a bluish appearance. Aluminum is only about one- third as heavy as iron. It is malleable, ductile, melts at 660, and is a good conductor of electricity and also of heat. It is used for wires and electric cables, for cooking utensils, and many other useful articles. Magnalium is an alloy consisting of 75 to 90 per cent of aluminum and 25 to 10 per cent magnesium. It is harder and lighter than aluminum and may be polished to a higher degree. Aluminum bronze consists of 90 to 95 per cent copper and 10 to 5 per cent aluminum. It is yellow in color, hard and strong, and is consequently often used in the arts. Aluminum paint consists of finely divided aluminum sus- pended in a suitable varnish or laquer. Mixed with the black oxide of iron, Fe 3 O4, finely divided metallic aluminum forms a mixture called thermite. When once this mixture is ignited (which can be accomplished by means of a fuse of magnesium ribbon to which is attached a mixture of magnesium powder and potassium chlorate) the chemical action continues rapidly and vigorously, the temperature developed being about 3000 C. The change that takes place is that aluminum robs the oxide of iron of its oxygen and so forms alumina and metallic iron, thus : 3Fe 3 O 4 + 8A1 = 4 A1 2 O 3 + 9 Fe black iron oxide aluminum alumina iron Thermite is used for welding car rails and making welds in ALUMINUM, HEAVY METALS, ALLOYS 143 many parts of machinery of various kinds. The parts need only to be butted together, a mold built around the joint, and the molten thermite mixture run upon it. Thus the iron is heated to the welding point very quickly. In many cases FIG. 47. Welding a car rail with thermite. it is not even necessary to take the machinery apart to make the repairs, the process is so simple. Aluminum sulphate, A1 2 (SO 4 ) 3 , is made by dissolving alu- minum hydroxide, A1(OH) 3 , in sulphuric acid, thus : 2 A1(OH) 3 + 3 H 2 SO 4 = A1 2 (SO 4 ) 3 + 6 H 2 O This salt is readily soluble in water and is used as a mordant for fixing dyestuffs upon fabrics. In the paper industry it serves in sizing paper. Ordinary potassium alum is a double sulphate of aluminum and potassium. It has the composition K 2 SO 4 . A1 2 (SO 4 ) 3 . 24 H 2 O. Ammonium alum and sodium alum contain am- monium and sodium respectively instead of potassium, otherwise they are quite like potassium alum. Alum crystal- 144 CHEMISTRY AND DAILY LIFE lizes in octahedra, dissolves readily in water, and like all other soluble aluminum salts, it has an astringent effect on the mucous membranes. Alum is used as a mouth wash in medi- cine, as a mordant in dyeing fabrics, and unfortunately it is also still used in baking powders. Alum baking powders contain alum or aluminum sulphate and baking soda. When this mixture is moistened, as, for example, in baking, carbon dioxide is evolved which raises the dough. There is simul- taneously formed sodium sulphate and aluminum hydroxide, and these remain in the biscuit, bread, or cake. The alumi- num hydroxide is readily dissolved by the hydrochloric acid in the stomach, thus forming aluminum chloride, a powerful astringent which interferes with digestion and is harmful to the linings of the digestive tract. Bread and cake made with alum baking powder looks well, for the carbon dioxide is liberated steadily and just about at the right rate to make fine-appearing cookery. As alum baking powders are cheap they -are still unfortunately much in use, even though in some states the law requires the fact that they contain alum or aluminum sulphate to be stated on the label to warn the public. It is a notable fact that though aluminum silicates are present in every soil, plants and animals only contain minute traces of aluminum. Clay consists essentially of aluminum silicates that have been formed from the weathering of feldspars and other complex silicates that make up the granitic rocks. In its purest form clay is white and is called kaolin, H 2 Al 2 (SiO 4 )2 . 2 H 2 O. This is used for making white porcelain ware. Common clay is discolored by impurities. Usually it is brown or red, which is caused by the presence of oxides of iron. Ordinary clay serves for making bricks, flowerpots, and the inferior and cheaper grades of pottery and crockery ware. ALUMINUM, HEAVY METALS, ALLOYS 145 In making bricks and porous earthenware the clay is molded into the desired shapes and then heated to redness in kilns, or " fired " as it is called. A cheap glaze may be put upon such ware by simply throwing salt into the kiln. Thus as the water is baked out of the clay, the water vapor at the high temperature that obtains decomposes the salt, forming hydro- chloric acid and sodium hydroxide. The latter unites with the surface of the clay, forming an easily fusible silicate which FIG. 48. A pottery kiln. Setting the shapes and getting ready to fire. produces a thin glassy coating or so-called glaze. Butter jars and other " stoneware " crocks are glazed in this way. Porcelain is vitreous throughout. This is accomplished by mixing finely pulverized feldspar and quartz with the kaolin before molding the dishes. When these are dried and finally fired, the feldspar and quartz fuse and fill the pores of the ware so that it exhibits a perfectly vitreous fracture instead of an earthy one, as the cheaper earthenware, either glazed or 146 CHEMISTRY AND DAILY LIFE unglazed, always does. Fire bricks contain a larger amount of silica than ordinary ones and are consequently more refractory. Fire clay also is especially rich in silica. Colored porcelain and colored glazes are produced by means of various metallic oxides, as in making colored glass, for example. Ultramarine is a fine dark blue pigment made by heating together clay, soda, sulphur, and charcoal out of contact with the air. It is probably a double compound of sodium alu- minum silicate and sodium sulphide. Ultramarine is attacked by acids, which destroy its color, but in presence of alkalies it is quite stable. It serves as laundry blue, also as a pigment in paints, and in the manufacture of wall paper. Further- more, it is used to destroy the yellow appearance of sugar, linen, cotton, and paper pulp, only enough being used to pro- duce a pure white appearance. Aluminum oxide, being a white powder that is practically insoluble in water and neutral in reaction toward litmus, is a typical " earthy " oxide, and aluminum is therefore termed an earth metal. There are several other similar earth metals, but they are so very rare and have no importance in practice that they need not be considered here. The oxides of cerium and thorium belong to these rare earths. It has already been stated that these are used in making mantles for the Welsbach lamps. In the Nernst lamp a filament of earthy oxides is heated to incandescence by passing the electric current through it. Platinum, gold, and silver are called the noble metals. They occur in nature in the free or uncombined state and are not at all readily oxidized. Platinum is silver-white, very malleable, and ductile. It occurs in alluvial sands in the Ural Mountains, in California, Brazil, and Australia. It is not attacked by hydrochloric, nitric, or sulphuric acid, nor by molten sodium or potassium ALUMINUM, HEAVY METALS, ALLOYS 147 carbonate, though aqua regia, a mixture of nitric and hydro- chloric acid, will dissolve it, forming platinic chloride. Plat- inum does not melt readily, its melting point being about 1777. These facts make it very valuable as a material out of which to construct dishes for certain kinds of chemical work in which glass and porcelain would be attacked. Platinum is also used for electri- cal contacts and for spark points in the spark plugs of gasoline engines. For the latter purpose it is commonly alloyed with indium, a metal which is similar to platinum. This alloy is the so-called hard platinum. Plati- num salts are used in photography, and also in chemistry for the purpose of estimating potassium, for potassium platinic chloride, K 2 PtCl 6 , is a beautiful, crystalline, yellow salt which is difficultly soluble in dilute alcohol. It is one of the very few potassium salts F I G 4 9 A . modern spark that do not dissolve copiously in water. In p i u g with a finely divided form platinum absorbs oxy- gen which is readily given off to oxidizable sub- stances. So, for example, one can light a gas jet by bringing it into contact with such finely divided plati- num, also called platinum sponge. Other metals of the platinum family are iridium, osmium, ruthenium, rho- dium, and palladium. They all occur with platinum in nature and exhibit somewhat similar properties. The many uses to which platinum has been put and the small supply make the metal very costly. It is worth about forty-five dollars per ounce. Gold, aurum, has always been regarded as an article of value. It melts at 1064, is extremely malleable and ductile. In aqua regia it dissolves, forming gold chloride. Gold coins 148 CHEMISTRY AND DAILY LIFE contain 9 parts gold and 1 part copper. This alloy is much harder than pure gold, which is rather soft and consequently will not wear well. For the same reason copper alloys of gold are used for jewelry, ornaments, etc. Pure gold is 24 carats fine ; 16 carat gold contains 16 parts gold and 8 parts copper; 14 carat gold contains 14 parts gold and 10 parts copper; etc. The LTnited States produces about 160 tons of gold annually, valued at about $96,500,000. The whole world produces about four times this amount per year. Gold salts are used in photography as " toning baths." Many metallic articles are plated with gold. This is accom- plished by immersing the article to be plated in a bath con- sisting of a solution of gold potassium cyanide, KAu(CN) 4 , together with an electrode of pure gold ; a current of electric- ity is then passed from the latter electrode through the solu- tion to the article to be plated, which is readily coated with the metal. Gold is also used on the edges and backs of books, for ornamenting chinaware, gilding signs, church spires, etc. Silver, argentum, occurs in nature in the free state, but also as the sulphide and chloride. It is much more abundant than platinum and gold. Nearly 2000 tons of silver are produced annually in the United States. The metal melts at 962, is malleable, ductile, and the best conductor of electricity and heat known. Sterling silver and silver coins contain 1 part cop- per and 9 parts silver. Nitric acid dissolves silver readily, forming silver nitrate, AgNOs, which is the most important of the silver compounds, for from it all others are prepared. It is used in medicine and in photography. It is one of the most soluble of all salts. Even in ice-cold water it is possible to dissolve 122 parts in 100 parts of water, while in boiling water the solubility increases tenfold. This salt is also called lunar caustic and is used for cauterizing wounds, removing warts, etc. Silver chloride, AgCl, silver bromide, AgBr, and ALUMINUM, HEAVY METALS, ALLOYS 149 silver iodide, Agl, are all insoluble in water. These salts darken on exposure to light, probably because of separation of finely divided silver. This fact is the basis of the use of these salts in photography. The photographic plate consists (B) FIG. 50. Photography. (A) A positive. (B) A negative. tially of silver bromide embedded in gelatine, which covers one side of the glass. On exposure to light in the camera, the bromide is slightly, but not visibly, reduced. When the plate is afterwards treated with a developer (a reducing agent like pyrogallic acid, hydrochinone, etc.), the reduction continues where the light has started it, and so the picture becomes visible. The developer must be washed off when the 150 CHEMISTRY AND DAILY LIFE picture has been sharply developed or the process will continue and produce a blurred outline. To remove the undecom- posed silver bromide in the gelatine, the plate (before expo- sure to light) must be treated with a solution of sodium thiosulphate, Na^Os, the " hypo " bath. This dissolves the silver bromide, leaving only the gelatine with the metallic silver particles which form .the outline of the picture. The plate is then thoroughly rinsed off and dried. It is a nega- tive, i.e. it is dark where the object which was photographed was light, and vice versa. To make a positive this negative is laid upon a bromide print paper, which is paper coated with a sensitive film of silver bromide similar to that on the bro- mide plate. The whole is now exposed to the light. The print must be " fixed " in the " hypo " bath the same as the negative, and it may then be exposed to the light. Silver plating is done in a solution of potassium silver cyanide, KAg(CN) 2 . The process is similar in all respects to gold plating, which has already been mentioned. The black stains that form on silverware (especially on spoons, forks, etc., which have been in contact with eggs) are silver sulphide, and not silver oxide as is often thought. Copper, cuprum, occurs in large quantities in the free state near Lake Superior. In Montana it is found in ores in combination with iron and sulphur. Nearly half of the an- nual output of copper in the world is produced in the United 'States, which contributes about 540,000 tons per year. Copper and gold are the only colored metals known. Copper melts at 1084. It is tough, rather hard, ductile, and malleable. It conducts heat and electricity well. Much copper wire is used in electrical work. For other purposes, however, copper is generally used in the form of alloys. These are made by melting the metals together in the desired proportions. The following are some of the most important alloys of ALUMINUM, HEAVY METALS. ALLOYS 151 FIG. 51. Plating silver spoons. The upper figure shows the rack on which the spoons are hung. 152 CHEMISTRY AND DAILY LIFE copper : Brass, which is yellow in color, contains 1 part zinc and 2 parts copper, though other proportions are often used. Dutch metal, which is reddish brown, consists of 1 part zinc and 5 parts copper. German silver consists of 80 to 95 per cent brass and 5 to 10 per cent nickel. Gun metal contains 9 parts copper in 1 part tin. Bell metal consists of 3 parts copper plus 1 part tin. The alloys of copper and tin are called bronzes. Bronzes for statuary com- monly contain 3 to 8 parts tin, 1 to 3 parts lead, 1 to 10 parts zinc, and 80 to 90 parts copper. Phosphor bronze is used in making parts of machinery. It is especially hard. It consists of bronze to which from 0.5 to 3 per cent phosphorus has been added. Copper dissolves readily in nitric acid, also in hot concentrated sulphuric acid, but cold sulphuric or hydrochloric acid has but very little effect on it. However, in presence of the air, many dilute acids do very gradually attack copper, which fact must be kept in mind, for copper and brass are frequently used for cooking utensils, and copper compounds are poisonom. On copper roofs, old coins, etc., that have been exposed to moist air for a long time there is formed a green deposit called verdigris. It is a basic carbon- ate of copper, CuCO 3 . Cu(OH) 2 . In ammonia water copper is soluble when in contact with the oxygen of the air. However, ammonia acts but slowly, hence it is often used for cleaning copper. The most important as well as the most common salt of copper is copper sulphate, blue vitriol, CuSCX . 5 H 2 O. It forms large blue crystals, is soluble in about 3 parts of water, and its solutions are used as a bath for copper plating, for spraying plants, especially in making Bordeaux mixture, and also for preparing other copper compounds. In Paris green copper also is present, as has been mentioned before. Mercury, quicksilver, hydrargyrum, is the only metal which is a liquid at ordinary temperatures. It is found in ALUMINUM, HEAVY METALS, ALLOYS 153 nature in the free state, but generally it is combined with cinnabar, HgS, which is red in color and is used as a pigment, being called vermilion. Mercury comes from Spain, Austria, Prussia, California, Japan, and China. It is used in ther- mometers, barometers, in amalgams for the backs of mirrors, also for filling teeth. Its compounds are used in medicine. So mercurous chloride, HgCl, is calomel. It is but slight'y soluble in water. Mercuric chloride, HgCl 2 , is corrosive sublimate. It is very copiously soluble in water and is an exceedingly powerful poison. It is to be borne in mind, more- over, that all mercury compounds are poisonous. The antidote is raw eggs and milk. The albumin forms an insoluble com- pound with the mercury salt, which is then got rid of by an emetic or a purge. Mercury melts at 39.4 C. and boils at 357 C. Solid mercury may be hammered into sheets and cut with tools like other metals. The alloys of mercury with other metals are called amalgams. The amalgam used in filling teeth commonly consists of tin, silver, and mercury. Gold alloys very readily with mercury. Indeed, the latter readily dissolves gold, and so is used in extracting fine particles of gold from the sands in which they occur. Mercury is also similarly used in silver mining. Tin, stannum, occurs in nature as tinstone, cassiterite, SnO 2 , which is found in England, Germany, Peru, Australia, Banca, and Alaska. To obtain the tin from this ore, the latter is heated with carbon, thus : SnO 2 + 2 C = 2 CO + Sn. Tin is very malleable, melts at 232, and boils at 1600. On exposure to the air tin remains nearly unchanged. It is used in making tin foil and tin plate. The latter consists of sheet iron coated with tin by the process of dipping the thoroughly cleaned iron into a bath of molten tin. Frequently copper is also treated by this method. Metallic tin is also used in 154 CHEMISTRY AND DAILY LIFE many alloys. Thus solder commonly consists of 1 part tin and 1 part lead, though other proportions are also employed. The bronzes, already mentioned, are alloys of tin and copper. Pewter consists of 1 part lead and 3 parts tin. Britannia metal contains 90 per cent tin, 8 per cent antimony, and 2 per cent copper. Among the most important salts of tin are stannous chloride, SnCl 2 , formed by dissolving tin in hot concentrated hydrochloric acid solution, and stannic chloride, SnCl 4 , produced by treating tin or stannous chloride with chlorine. Stannous chloride forms white crystals that are very soluble in water, while stannic chloride is a fuming liquid which boils at 114. With ammonium chloride stannic chloride forms a double salt, SnCl 4 . NH 4 C1, called pink salt, which is used as a mordanto Stannous chloride is a reducing agent, for it readily passes over into stannic chloride, for example ; SnCl 2 + HgCl 2 = SnCl 4 + Hg stannous mercuric stannic mercury chloride chloride chloride Mosaic gold is stannic sulphide, SnS 2 , prepared by heating together sulphur and tin. It consists of golden yellow crys- tals, and is used for " bronzing " articles. Tin is a rather costly metal. The world produces about 110,000 tons of it annually. Most of this is mined in Banca and other neigh- boring East India islands. Lead, plumbum, is soft and malleable. It melts at 327 C. and it is 11.4 times heavier than water. This metal has been used in the arts even in ancient times. The Romans used lead for water pipes, and it serves for this purpose to this day. The chief ore of lead is the sulphide PbS, which crystal- lizes in cubes. It is commonly known as galenite. As already mentioned, solder and pewter are alloys of lead and tin. Shot and bullets consist of lead containing about three- ALUMINUM, HEAVY METALS, ALLOYS 155 tenths per cent arsenic, while in Babbitt metal from 70 to 90 per cent lead is alloyed with antimony and tin. Much lead is used also in the manufacture of lead storage batteries and in making lead salts and other compounds. The lead storage battery consists essentially of a lead plate and a lead plate coated with lead peroxide, PbO 2 , dipping in a solution of sul- phuric acid of specific gravity 1.2. As the battery acts, the current flows from the lead through the solution to the lead peroxide. Thus lead is dissolved and hydro- gen is set free at the other plate, but it at once attacks the lead peroxide, forming lead oxide and water. The lead oxide formed is dissolved by the sulphuric acid present, forming water and lead sulphate. The electromotive force of the battery is two volts. Charg- ing the battery consists simply in conducting a current from a dynamo through the battery from the peroxide plate through the solution to the lead plate. Thus lead is again deposited on the lead plate and lead per- oxide is again formed on the opposite plate ; that is to say, the chemical changes that took place when the battery was discharged are reversed. These batteries are used for running electric automobiles, also for lighting and ignition purposes, etc., for they have a high electromotive force and a low in- ternal resistance and so yield a powerful current. Litharge, lead oxide, PbO, is a yellow powder which is used extensively in glazing pottery and ironware, in putting decora- tions upon porcelain, and in making glass that has a high in- FIG. 52. A lead storage battery. 156 CHEMISTRY AND DAILY LIFE dex of refraction. It is also employed in making various lead salts. The most common soluble lead salts are the nitrate, Pb(NO 3 ) 2 , and the acetate, Pb(C 2 H 3 O 2 ) 2 . 3 H 2 O. The latter is commonly known as sugar of lead, for it has a sweetish, yet dis- agreeable taste. The basic lead acetate, Pb(C 2 H 3 O 2 ) 2 . (PbO),, is often used as a lotion for diseased parts, a 2 per cent solution of this salt being called lead water. Lead arsenate, Pb 3 (AsO 4 ) 2 , is used for poisoning potato bugs and also for spraying trees and shrubs. Like Paris green, lead arsenate is but sparingly soluble in water. White lead is a basic car- bonate of lead, Pb(OH) 2 . 2 PbCO 3 . It is sold in large quan- tities ground in linseed oil. On being further diluted with the latter, this mixture makes an excellent paint, especially for wood that is to stand exposure to the weather. White lead paint is not infrequently adulterated with chalk, so-called whiting, barium sulphate, or lead sulphate. It is to be borne in mind that all lead compounds are poisonous. The more soluble the compounds are, the greater is the danger of being poisoned by means of them. Painters are often poisoned by the white lead they use. This gets on their hands, and it is very likely that when these touch the nose and lips, the lead compounds come into contact with the mucous mem- branes and are absorbed. Thus small quantities are taken into the system at a time, but these accumulate and finally cause lead colic. Chromium is a hard, steel-gray, brittle metal which melts at about 1515 C. It is 6.8 times as heavy as water. It occurs in chrome iron ore, Cr 2 O 3 . FeO, and is used in making chrome steel, which is steel alloyed with a small percentage of chromium. This makes a very hard steel. Chromic oxide, Cr 2 O 3 , is chrome green and is used as a pigment in paint. It is also used in making green glass and glazes. Chrome yellow is lead chromate, PbCrO 4 . It is also used as a ALUMINUM, HEAVY METALS, ALLOYS 157 pigment in paint. The most common soluble compound of chromium is potassium bichromate, K 2 Cr 2 O7. It forms beautiful orange-colored crystals. It semes as an oxidizing agent, also in chrome tanning and in dyeing fabrics. Tungsten is also a brittle metal. In some of its properties it resembles chromium. Tungsten is used in making tungsten steel, and also in producing the filaments for the incandescent tungsten electric lamps. Molybdenum is closely allied to tungsten. Ammonium molybdate, (NH 4 ) 2 MoO4, is used in testing for the presence of phosphates, for when added to a nitric acid solution of the lat- ter there forms a very characteristic yellow precipitate called ammonium phosphomolybdate, (NH^PCX . 11 MoO 3 . 6 H 2 O. This precipitate is soluble in ammonia water. Zinc is about 6.9 times as heavy as water, melts at 420 C. and boils at 918 C. It is found in nature mainly as the carbonate ZnCO 3 , and the sulphide ZnS. The latter ore is popularly known as blackjack. These ores are first heated in contact with the air and are thus changed to zinc oxide. ZnCO 3 on heating = ZnO + CO 2 zinc carbonate zinc oxide carbon dioxide ZnS + 30 = ZnO + SO 2 zinc sulphide oxygen zinc oxide sulphur dioxide (from the air) From its oxide, zinc is then obtained by heating with carbon in earthenware retorts, thus : ZnO + C = Zn + CO zinc oxide carbon zinc ] carbon monoxide In the form of sheets zinc is much used for gutters, roofs, orna- ments, etc., on buildings. Galvanized iron, so called, consists of sheet iron which has been coated with zinc by the process of dipping the thoroughly clean iron into a bath of molten 158 CHEMISTRY AND DAILY LIFE FIG. 53. A sal ammoniac battery, showing the parts. zinc. The beautiful flaky appearance of galvanized iron is due to the crystals of zinc on its surface. The coating pro- tects the iron so that it will not rust. Zinc is also used in many primary electrical batteries. In the common battery used for ringing doorbells, Fig. 53, there is a solution of ammo- nium chloride, sal ammoniac, into which dip a stick of zinc and' a cylindrical plate of carbon. The latter is usu- ally so shaped as to form also the cover of the jar. The battery carbon is made by grinding coke with black strap molasses or coal tar as a binder, molding, and then heating the product at first gently, and finally to redness out of contact with the air. In this battery, the current flows from the zinc through the solution to the carbon. Zinc chloride is thus formed as the zinc and hydrogen is evolved at the carbon. The hydrogen is absorbed by the carbon and finally escapes in the air. Dry batteries contain the same in- gredients. Only here the outer casing is made of zinc, which also serves as the negative pole of the battery. The carbon generally is a cylin- drical piece in the middle of the cell and is sur- rounded by manganese dioxide which oxidizes the hydrogen that is liberated. Plaster of Paris is used to absorb the liquid and hold the ingredients in place. But the contents must be moist or no current can flow. Blue cup bat- teries, Fig. 55, contain a zinc pole in dilute sulphuric acid, and a copper pole in saturated copper sul- phate solution. As the current passes from zinc through the FIG. 54. A dry battery. ALUMINUM, HEAVY METALS, ALLOYS 159 solution to the copper, zinc dissolves from the zinc plate, and copper deposits on the copper plate. Other metals could be used in place of zinc, but the latter is on the whole the most economical, pro- ducing a relatively high electromotive force. Brass is the most impor- tant alloy of zinc which is in common use, but the latter metal is often present in riG . 55 ._A blue cup battery and its parts. other alloys. About two hundred and fifty thousand tons of zinc are produced in the United States each year, and the rest of the world produces about two times that amount. The compounds of zinc are analogous to those of magnesium. It is to be kept in mind, however, that zinc compounds are poisonous. Zinc oxide, ZnO, is used in paints, being called zinc white. Mixed with lard or vaseline, it forms zinc oxide ointment. Zinc chloride, ZnCl2, is very soluble in water. It is used for preserving wood, especially railroad ties. The salt permeates the wood and is an antiseptic ; thus it is death to the germs that are the cause of the decay of wood. White vitriol or zinc sulphate, ZnSO 4 . 7 H^O, is another very common zinc salt. It is readily soluble in water. Manganese is a hard, brittle metal, which is eight times as heavy as water, and melts at 1300. In na- ture it is found mainly as pyrolusite, which is manga- nese dioxide, MnO 2 . Alloys of manganese and iron are used injhe steel industry. With copper, manganese forms alloys called manganese bronze. These are very hard and strong. Manganese dioxide is often used in preparing chlorine, thus : 160 CHEMISTRY AND DAILY LIFE MnO 2 + 4HC1 = 2H 2 O -f MnCl 2 + C1 2 manganese hydrochloric water manganous chlorine dioxide acid chloride The salts of manganese are quite numerous. The manganous salts are pink in color, and the chloride, MnCl 2 , sulphate, MnS0 4 , nitrate, Mn(NO 3 )2, and acetate, Mn(C 2 H 3 O 2 ) 2 , are soluble in water. Manganese acts as an acidic element in manganates and permanganates. Of these salts potassium permanganate, KMnO 4 , is of practical and commercial impor- tance. This salt crystallizes in needles that are of a dark purple, lustrous hue. The salt is soluble in water and very often serves in the laboratory as an oxidizing agent. It is also used as a disinfectant. Nickel is malleable, ductile, and 8.9 times as heavy as water. Its melting point is about 1485 C. Nickel coins consist of 25 per cent nickel and 75 per cent copper. When alloyed with brass, nickel forms German silver. Nickel steel is used for making armor plates for warships. Nickel does not tarnish readily on exposure to air, and hence it is often used to plate iron, copper, brass, etc. Nickel plating is carried on in the same way as gold or silver plat- ing. The bath for nickel plating consists of a solution of nickel ammonium sulphate, (NH 4 ) 2 SO4 . NiSO 4 . 6 H 2 O, and, of course, a plate of solid nickel is placed in the solution opposite to the object to be plated. The electric current is then passed from the nickel plate through the solution to the object that is to receive the nickel coating. The salts of nickel are green, and when soluble they yield green solutions. Cobalt is analogous to nickel in most important respects. Its salts when dissolved in water yield dark red solutions, and the crystals that deposit from such solutions are also red, for in general the crystals contain water of crystallization. When this water is driven off, anhydrous, i.e. dry, compounds ALUMINUM, HEAVY METALS, ALLOYS 161 are formed which are blue. Cobalt silicate is blue, and so glass containing cobalt silicate is blue glass, also called smalt glass. Indeed, about the only practical use which is at present made of cobalt is in the production of blue glass, blue porcelain, and blue glazes on enamel ware. Cobalt, nickel, and iron are metals that are attracted by a magnet, and are consequently said to be magnetic. Of all the metals, iron, ferrum, is the most use- ful. It very rarely is found in the free state in nature except in me- teorites. The most im- portant ore of iron is hematite, which is an oxide having the com- position Fe2O 3 . It is dark red in color and when finely ground it serves as a pigment in paint, being called red ocher. The rich deposits of iron ore in the Lake Superior re- gion consist of hematite. Magnetite, or magnetic iron ore, Fe 3 O 4 , is black. Limonite, 2 Fe 2 O 3 . 3 H 2 O, is a hydrous oxide of iron. It is yellow in color. It too serves as a pigment under the name of yellow ocher. These oxides and siderite, ferrous carbonate, FeCO 3 , form the chief ores of iron. FIG. 56. A blast furnace in which cast iron is made. 162 CHEMISTRY AND DAILY LIFE To obtain metallic iron from these ores, they are mixed with charcoal, coke, or coal and limestone and then heated in a blast FIG. 57. A modern blast furnace, showing the pig iron bed. furnace. In the lower part of the furnace carbon dioxide is formed, for here air is blown in and there is consequently sufficient oxygen to form carbon dioxide. The latter gas passes upward through the hot carbon and is converted into ALUMINUM, HEAVY METALS, ALLOYS 163 carbon monoxide, thus : CO 2 + C = 2 CO Then the carbon monoxide reduces the iron ore as follows : Fe 2 O 3 + 3 CO = 2 Fe + 3 CO 2 and out of the top of the furnace there issues carbon dioxide, mixed, however, with much carbon monoxide. This gas is now used to run gas engines that furnish the air blast for the furnace. The molten iron accumulates in the bottom of the furnace and is tapped off from time to time. The iron is run into molds of sand so as to form rough bars, or pigs as they are called. The cast iron thus obtained is known as pig iron. The limestone was added so that the sand and silicates might react with the lime to form a fusible calcium silicate which floats on top of the molten iron and is run off. It is the so-called slag. It is now used for making Portland cement, whereas formerly it was thrown away. Cast iron always contains considerable amounts of carbon and other impurities. In fact chemically pure iron is not obtainable in the market, being practically unknown. The different kinds of iron and steel, then, are really iron containing various other substances. The carbon content especially determines the properties of the iron. Cast iron usually contains from 2 to 3 per cent of carbon embedded in the iron FIG. 58. Cast iron as it appears under the microscope. 164 CHEMISTRY AND DAILY LIFE as graphite, and from 1 to 1.5 per cent of carbon as iron carbide, i.e. chemically combined carbon. From 0.5 to 4 per cent of silicon, from 0.4 to 2 per cent of phosphorus, and sulphur up to 0.2 per cent are also generally present. Wrought iron is the purest iron on the market. It contains usually less than 1 per cent of foreign ingredi- ents. The carbon con- tent is generally less than 0.2 per cent. Wrought iron melts at 1600 C. and can be welded at from 900 to 1100. In welding, borax or sand is used. Thus a slag of iron borate or iron silicate is formed, and when the iron, brought to welding heat, is pounded together, this slag flies off and the clean sur- faces of the pieces of iron come into intimate contact so that the forces of cohesion can act and hold them together. Wrought iron is made from cast iron by puddling, a process which consists of adding iron oxide to the pig iron and then heating in a current of air in a so-called reverberatory furnace. Thus the carbon of the pig iron is removed because it unites with the oxygen of the iron oxide and of the air. Other impurities present, like silicon and phosphorus, are also oxidized and form a slag with some of the iron. Finally the iron is almost free from foreign matter except a few tenths of a per cent of carbon. Malleable iron is a cheap sub- stitute for wrought iron. It is made by embedding the cast- ings in iron oxide and then heating for about two days, after FIG. 59. Wrought iron as it appears under the microscope. ALUMINUM, HEAVY METALS, ALLOYS 165 which all is cooled off slowly. Thus some of the carbon is removed from the castings, making them less brittle than they were. It is to be borne in mind that a large carbon content makes iron brittle. Sulphur and phosphorus are especially objectionable in iron, for they make it quite brittle. Steel is practically free from sulphur and phosphorus. Silicon too is present only in small amounts. The carbon in steel varies from about 0.2 to 1.6 per cent, mild steel containing the smaller amounts of car- bon. The process of hardening and temper- ing consists of heating the steel and then chill- ing it. Very sudden chilling does not give the carbon a chance to crystallize out as graph- ite, and the combined carbon makes the steel wry hard. If the cool- irig proceeds Very FlG . 60 ._ steel as it appears under the mi- slowly, practically all croscope. the carbon crystallizes out, and a soft material is obtained. By suitable chilling, the proper temper, that is, the desired degree of hardness, may be obtained. The removal of the carbon from cast iron so as to produce steel consists essentially in oxidizing the carbon. This is accomplished by blowing air into the molten mass, as in the Bessemer process, in which a so-called Bessemer converter is used, or by heating cast iron with iron oxide on the hearth of a furnace, the so- called open hearth process. When the cast iron contains much phosphorus, the hearth of the furnace is made of lime 166 CHEMISTRY AND DAILY LIFE obtained by calcining dolomite. This absorbs the phos- phorus, phosphates of calcium and magnesium being formed which are sold as fertilizer. Nearly thirty million tons of iron are produced yearly in the United States. FIG. 61. An open hearth furnace. Iron forms two series of salts, the ferrous and the ferric. Thus there are ferrous chloride, FeCl 2 , and ferric chloride, FeCl 3 , ferrous oxide, FeO, and ferric oxide, Fe 2 O 3 , etc. In the rusting of iron hydrated ferric oxide is formed. The black oxide, also known popularly as hammer black, which results when iron is heated in the air, is ferrous ferric oxide, FeO . Fe 2 O 3 or Fe 3 O 4 . It is magnetic. Iron nearly always acts as a basic element, and in this capacity it forms salts with practically all of the various acids. Only a few of the most important salts of iron will be mentioned here. Ferrous sulphate, also known as green vitriol or copperas, is FeS0 4 . 7 H 2 O. It is formed by the oxidation of pyrite, ALUMINUM, HEAVY METALS, ALLOYS 167 FeS 2 , which occurs in nature and is called fool's gold, for its crystals are of a lustrous, golden yellow appearance. Fer- rous sulphate dissolves readily in water. It is used in making ink. The latter contains tannin and ferrous sulphate. The tannin is an essential ingredient in extract of nutgalls, which is added to the ferrous sulphate solution. Ferrous sulphate is also used as a disinfectant, as a mordant in dyeing fabrics, and as a reducing agent. This salt is cheap and readily obtainable from dealers everywhere. Ferric chloride, FeCl 3 . 6 H 2 O, forms a dark brown crystal- line mass, not unlike maple sugar in appearance. It dissolves copiously in water, also in alcohol and ether. It is used in medicine. The so-called styptic cotton which is used to stop the bleeding of wounds consists of absorbent cotton treated with a solution of ferric chloride. This material not only allays bleeding by forming a clot with the blood, but it also acts as an antiseptic, thus protecting the wound from germs. Blue print paper consists of paper that has been coated with a solution of ferric ammonium citrate plus potassium ferric cyanide. The latter is also called red pmssiate of potash, K 3 Fe(CN) 6 . This paper is thus coated and dried in the dark. On exposure to light, the ferric citrate is in part reduced to ferrous citrate, and the latter, like other soluble ferrous salts, reacts with the potassium ferric cyanide, forming a blue pre- cipitate, Fe 3 " [Fe"' (CN) 6 ] 2 , Turnbull's blue. Where the paper has been protected, no precipitate forms ; and so when after exposure to light in the printing process the paper is washed with water, the parts on which the light has acted appear blue, whereas the parts that have been protected appear white, for the original coating on the blue print paper is soluble in water. It is to be remembered that if blue print paper, not exposed to light, is washed with water in the dark, white paper is obtained. The blue printing is done in frames 168 CHEMISTRY AND DAILY LIFE similar to those used for printing photographs from negatives. The latter will also serve for making blue prints. But blue prints are commonly made by using tracing cloth or paper that is not too thick on which the writing or drawing that is to be reproduced in blue print form has been traced. The thicker the paper, the longer must be the exposure to the light to get the desired results. Ammonia, caustic soda, and caustic potash solutions decompose TurnbuU's blue, and so may be used to write white characters on blue prints. The latter are stable in the light, also toward acids, but alkalies, as just mentioned, destroy the blue color. It is to be borne in mind that in small amounts iron is present almost everywhere. In all soils it occurs. To the brown sand, earth, and the brownish and yellowish clays it gives their characteristic colors. Our sandstones and other rocks all contain compounds of iron. Indeed, the grains of sand in the sandstones are commonly cemented together with oxides of iron. Iron compounds are present in the tissues of all plants and animals. The green leaf contains chlorophyl, for the formation of which iron is essential, and the blood of animals contains haemoglobin, in which iron plays an important role. In fact, without iron plants and animals cannot live. Nevertheless, it must be remembered that the quantity of iron present in living beings, exceedingly important though it is, is after all small ; so, for instance, the human body contains only about 0.004 per cent of iron. QUESTIONS 1. Where is aluminum found in nature, and in what form ? 2. How is metallic aluminum made ? 3. Mention the most important characteristics of metallic alu- minum. What is it used for ? 4. What is thermite ? Describe its use. ALUMINUM, HEAVY METALS, ALLOYS 169 5. What is alum, and what is it used for ? 6. How is porcelain made ? 7. What is the difference between porcelain and ordinary crockery ware ? 8. How is ultramarine made ? What are its uses ? 9. How much alumina can be made from 45 grams of potassium alum? 10. What are the noble metals ? Describe each. 11. For what purposes is gold used ? Why? 12. What is the most important compound of silver ? What are its uses ? 13. What is sterling silver ? What is it used for ? 14. Discuss the use of silver compounds in photography. 15. What is a negative ? 16. Where does copper occur in nature, and what are its chief characteristics and uses ? 17. Mention three important alloys of copper. 18. Describe the most common compound o"f copper. 19. What is Bordeaux mixture ? 20. Of what use is mercury ? 21. Mention two important compounds of mercury and state what they are used for. 22. How is tin obtained from its ores ? 23. What is pewter ? Britannia metal ? Solder ? 24. What is pink salt and what is it used for ? 26. What is mosaic gold ? State its use. 26. Mention the properties and uses of lead. 27. What does a lead storage battery consist of ? 28. What is lead arsenate? Sugar of lead? White lead? What is each used for ? 29. What use is made of potassium bichromate ? 30. What is ammonium molybdate used for ? 31. How does zinc occur in nature? How is it obtained from these ores ? 170 CHEMISTRY AND DAILY LIFE 32. What is galvanized iron ? 33. Describe an ordinary battery such as is used for ringing a doorbell. 34. State the uses of zinc oxide and zinc chloride. 35. What is manganese dioxide used for? Write the equation expressing the chemical changes. 36. Of what use is potassium permanganate ? 37. State the characteristics and uses of nickel. 38. How proceed to nickel plate an iron spoon ? 39. What use is made of cobalt ? 40. Name the chief iron ores and state briefly how iron is obtained from its ores. 41. What is the chief difference between cast iron and wrought iron? 42. How is steel made, and how is it tempered ? 43. Why are sulphur and phosphorus objectionable in iron? 44. What is malleable iron ? How is it made ? 45. Explain how fertilizers and also Portland cement are made in connection with the production of iron and steel. 46. How many series of salts does iron form ? Mention two com- pounds of each series. 47. What is ordinary ink ? 48. Describe how to make a blue print. 49. Discuss the occurrence of iron in plants and animals. CHAPTER XII PAINTS, OILS, AND VARNISHES THE use of paints, oils, varnishes, and lacquers dates back to ancient times, for the great advantage of putting a pro- tective coating upon articles of wood, metal, etc., has been appreciated for many centuries. That such coatings fre- quently enhance the beauty of the objects to which they are applied was also duly recognized. It has already been stated that oils may be divided into two great classes: (1) the mineral oils, and (2) the oils of organic, that is to say, of plant or animal, origin. But oils may also be classified according to their use. When an oil is to be used in paint it must be .an oil that will dry and form a hard coating. Oils like lard oil, olive oil, cottonseed oil, goose oil, neat's-foot oil, and the various mineral oils obtained from petroleum will not dry. They remain smeary when they are spread upon wood, glass, metal, etc., and they are conse- quently called non-drying oils. On the other hand, linseed oil and poppy-seed oil, for example, will, when similarly spread out upon objects, gradually dry to a hard, firm, resistant coating. Such oils are termed drying oils. While mineral oils are obtained from the distillation of crude petroleum, animal oils are obtained by heating animal fats and then subjecting them to pressure, and vegetable oils are prepared by grinding seeds and expressing the oil from them either at room temperature or at higher temperatures. Sometimes solvents like carbon bisulphide and gasoline are used in extracting animal or vegetable oils from fatty animal or plant 171 172 CHEMISTRY AND DAILY LIFE matter. Thus the fats pass into the solvent, and the clear solution containing the fat may be filtered from the solid residue that remains. From the solution the oil or fat is obtained by distilling off the sol- vent, which can be used over and over again. Linseed oil is the most impor- tant of all oils for use in paints. It is obtained from flaxseed, one bushel of the latter yielding about 2.3 gallons of oil. The dry- ing qualities of linseed oil depend upon the fact that it has the poiver to absorb oxygen from the air. This oxidized linseed oil, also called linoxin, forms a tough, hard, resistant film and is quite unlike the original oil. By boil- ing linseed oil with lead oxide, oxides of manganese, or linoleates of these metals, the oil dries more rapidly, and the substances used to bring about this effect are called driers. They are, in general, oxidizing agents whose purpose is to increase the rapidity of the oxidation of linseed oil to linoxin. Linseed oil itself is essentially a glycerine salt of linoleic acid, of the formula (CigHaiO^ . C 3 H 5 O 3 . There are also present in the oil, to the extent of about 20 per cent, olein, palmitin, etc. American linseed oil has a specific gravity of 0.9336 at 15 when raw, whereas the boiled has a specific gravity of 0.938. It boils at 130 C. The raw oil obtained by pressing ground flaxseed in the cold is very light- colored, whereas oil that is pressed from the hot seeds, and also boiled oil, is dark brown in color and has a greenish tinge. FIG. 62. A press used in re- moving oil from seeds. PAINTS, OILS, AND VARNISHES 173 Upwards of fifteen million bushels of flaxseed are raised annually in the United States, which would represent about thirty-four million gallons of linseed oil. Being somewhat expensive, linseed oil is often shamefully adulterated with cheaper oils, among which may be mentioned fish oil, petro- leum oils, rape, cotton, hemp, and corn oil. The presence of these adulterants, of course, lowers the quality of the paint made by using the admixture. Rosin is also frequently used to adulterate linseed oil. When linseed oil is thoroughly mixed with white lead, an excel- lent paint is obtained. This is usually slightly yellowish in tint, due to the color of the oil, but by adding a very small amount of blue Prussian blue, for example this yellow shade is dispelled and a pure white paint is obtained, which FIG. 63. A barrel of white lead. dries rather slowly, but is exceedingly durable, especially for outside work that is to be exposed to the weather. The paint can be made to dry more rapidly by adding small amounts of so-called driers, sometimes also termed Japan driers. As stated before, they consist of substances that absorb oxygen from the air and then give it off to the oil, at 174 CHEMISTRY AND DAILY LIFE least in part, thus hastening the oxidation, that is, the drying, of the oil. Turpentine is a substance of this kind, and so its addition to paint hastens the process of drying. Driers are commonly made by heating lead and manganese oxides (about four pounds of the mixture) with a gallon of linseed oil to about 500 F. and stirring till a uniform mass results. This is then dissolved in turpentine before it is cold. By adding some of the solution thus obtained to paint, the latter dries rapidly. Too much drier must not be used, or the paint after hardening will keep on oxidizing, and, thus the oil will be oxidized too much and spoiled, or, as the painters say, " burnt." Any drier really tends to make the paint less durable. There is much linseed oil sold as boiled oil which is not boiled at all, being raw oil to which drier has been added; it is popularly called " bunghole boiled oil." Often zinc oxide is used in paints together with white lead, even to the extent of half and half. Paints may be colored by adding suit- able pigments. These should be finely ground and then thoroughly mixed with the white lead and oil. Stirring is really not as good as grinding all together. The pigments remain suspended in the form of minute particles in the paint and also in the film as it dries after it has been applied. As already stated, yellow ocher and red ocher are oxides of iron. They are cheap, permanent, and excellent pigments. Chrome yellow gives a beautiful canary-yellow shade, but it is more costly. Vermilion is suphide of mercury, cinnabar ; and chrome green, which is very brilliant in hue, is chromic oxide. Prussian blue is produced by adding ferric chloride to a solution of potassium ferrocyanide, K 4 Fe(CN) 6 . It is often used as a pigment. Ultramarine blue too is used, but it is not as permanent. It is excellent for interior work in kalsomines, etc. Shades of gray result by adding small quantities of lampblack to white paint. // is always best PAINTS, OILS, AND VARNISHES 175 to apply two or more thin coats of paint ivell brushed out, rather than to attempt to cover the object with but one thick coat. This is obvious, for a thick coat will not dry readily, and will tend to peel off. As adulterants of white lead, chalk, barytes, lead sulphate, kaolin, and a mixture of zinc sulphide and barytes, which is termed lithophone, are frequently used. None of these should be present in high grade paint. Varnish is made by care/idly melting rosin and then stirring in linseed oil and heating the mixture together. This is finally FIG. 64. Gathering turpentine. thinned to the proper consistency with turpentine. Tur- pentine, CioHie, is obtained by distilling the sap collected from pine trees. After the turpentine has been distilled off, 176 CHEMISTRY AND DAILY LIFE the rosin remains behind in the retort. Many paints contain varnish, which gives a gloss to the film that forms on drying. Sometimes benzine or gasoline are used to thin varnishes or paints instead of turpentine, but this is bad practice, for the film that remains after the benzine has evaporated is not durable. The turpentine never wholly evaporates ; it always leaves a film of its own which together with the varnish or paint film makes for greater durability. If, instead of rosin, copal, a fossil rosin found in Africa, is used in making varnish, a far better article is obtained, but it is also much more costly. White enamel paints are superior varnish to which white lead or zinc oxide or both have been added as a pig- ment. Other colored enamel paints result by introducing suitable pigments. Shellac varnish consists of a solution of shellac, a plant resin, in alcohol. Either wood alcohol or grain alcohol may be used. White shellac is bleached shellac, whereas the brown or golden shellac has the natural color of the resin. Shellac varnish is cheaper, but not as durable, as varnish made from oil and rosin. FIG. 65. Trinidad asphalt lake near the edge. PAINTS, OILS, AND VARNISHES 177 Black varnishes used for coating iron are commonly pre- pared by dissolving coal-tar pitch or asphaltum in turpentine or benzine or a mixture of the two. Asphalt, also called asphaltum, is a natural pitch. Large quantities of it are found in Trinidad. It is also used for making excellent street pavements, by mixing it with sand, dust, and finely crushed stone and then applying this on a foundation of con- crete, rolling the hot mixture securely into place by means of heavy, hot steam rollers. FIG. 66. An asphalt street in New York. Kalsomines consist of a thin solution of glue to which slaked lime or chalk or both have been added. These mix- tures may be tinted by means of pigments just as in the case of paints. Many kalsomines contain plaster of Paris, which forms a hard coating as it dries. Kalsomines, being soluble in water, can be used only for indoor work, preferably on plastered walls. 178 CHEMISTRY AND DAILY LIFE QUESTIONS 1. For what purposes are paints and varnishes used? 2. How may oils be classified ? 3. What is a drying oil ? Name one and state why it dries. 4. How are plant oils obtained ? Name several plant oils. 6. How are animal oils prepared? Name several. Are these suitable for use in paint ? Why ? 6. What is a so-called drier ? How are driers used ? 7. What are some of the adulterants used in linseed oil ? 8. Why is turpentine better than benzine in paints and var- nishes ? 9. Describe how to make a good paint for outside woodwork. 10. Mention five pigments that are used in paints. 11. How is varnish made ? 12. What is shellac varnish ? 13. What is black varnish and for what purpose is it used ? 14. What is asphalt ? For what is it used ? 15. What is kalsomine ? CHAPTER XIII LEATHER, SILK, WOOL, COTTON, AND RUBBER WEARING apparel is made of materials derived from animals or plants. Shoes are made of leather, which is produced FIG. 67. Silkworm, cocoon, egg-mass, and strands of silk. 179 180 CHEMISTRY AND DAILY LIFE from the skin of animals. Furs, too, are skins of animals that have been subjected to certain treatment. Silk is made from the cocoon of the silkworm. The silk threads are spun and then woven into cloth. Wool is the hair of the sheep. It is cleansed, freed from fat, spun, and then woven. Wool and silk are substances containing carbon, hydrogen, oxygen, sulphur, and FIG. 68. A sheep. nitrogen. They are very like horn, hoofs, and finger nails in chemical composition, and when burnt all of these give a pecul- iar odor, namely, the odor of burnt hair. This odor is due to the nitrogenous decomposition products that are formed. Thus it is quite easy to distinguish between wool and silk on the one hand, and cotton on the other, for the latter does not evolve the odor of burnt hair when ignited. Cotton, as has already been pointed out, is practically pure cellulose (C 6 HioO 6 ) OT and so contains no nitrogen. Moreover, wool LEATHER, SILK, WOOL, COTTON, RUBBER 181 and silk are readily soluble in caustic alkalies, whereas cotton is not. Cotton, too, is freed from mechanical impurities, and then spun and woven. Because of the great chemical differences between cotton and wool, they react somewhat differ- ently toward dyestuffs ; and if a fabric containing both cotton and woolen threads is dyed in the same liquid, their threads will appear of different shades. Frequently cloth is made by using cotton warp and woolen woof, and the rather rough bulky nature of the woolen threads after the process of milling FIG. 69. A fleece of wool. hides the cotton threads. Such cloth is often quite service- able, though it is much cheaper than cloth that is all wool, for cotton is far cheaper than wool. Old woolen clothing is frequently taken apart and used over again to make new cloth. In this process the threads of wool are necessarily much shortened and the cloth woven therefrom is much less durable than new goods. By using some cotton in the warp of such made over woolens, a better wearing cloth is obtained. When used over again and again, however, the threads finally become so short as to be worthless except as a fertilizer. 182 CHEMISTRY AND DAILY LIFE Moreover, when used for a second time, there is cotton mixed with the woolen threads, and such material woven on cotton warp is commonly known as shoddy. Silk, woolen, cotton, and linen fabrics are now dyed almost exclusively by means of FIG. 70. An old-fashioned loom of colonial days. the aniline dyestuffs, also called coal-tar colors. As a rule the colors are more readily fixed upon silk and woolen fibers, and less readily upon the cotton and linen fibers. Linen is made of flax. Like cotton, it is essentially cellulose. Mordants often have to be used to fix dyestuffs on cotton and linen, and sometimes also upon wool and silk. Dyestuffs are in general either slightly acid or basic in character. Mordants, too, are LEATHER, SILK, WOOL, COTTON, RUBBER 183 either weakly acidic or basic substances. An illustration will serve to show the use of a mordant. Suppose the mor- dant is tannic acid ; by first dipping the fabric into a solu- tion of this acid, the fibers take up a certain amount of the substance ; and when they are now immersed in a dye of basic properties, the latter is fixed upon the fiber. In making leather, the object is to produce a material that is strong, pliable, insoluble in water, and not subject to putrefac- tion. To this end the hides or skins are first thoroughly softened by soaking them in water. They are then immersed in a bath of slaked lime (milk of lime), which loosens the hair so that it can afterward be readily removed mechanically. The process of removing the hair is called depilation. It is also sometimes accomplished by simply hanging the moist skins up so that slight putrefaction ensues, whereupon the hair becomes loose and can readily be removed. If the lime process of depilation has been employed, the lime in the skin must be removed before proceeding farther. This is done by treating in a bath of dilute sulphuric acid and then washing with water. This treatment greatly increases the bulk of the skin, and thus it is in excellent condition for the next step. This consists of immersion in extract of hemlock or oak bark. This extract contains tannin which unites chemically with the fibers of the hide and produces the compound which is called leather. Leather does not dissolve in water and will not putrefy. It requires about seventy days to convert the hide into leather by immersing the hide in extract of tanbark as described. Often hides are put into vats with alternate layers of ground tanbark, and water is then added to cover all. Thus the water extracts tannin from the bark, and leather is formed by the union of the hide with the tannin. This process requires a much longer time. Two or three months, or even a year, may be required to secure the end in 184 CHEMISTRY AND DAILY LIFE view. Moreover, the bark becomes exhausted and the hides have to be placed in another vat with new bark from time to time. Nevertheless, in this way an excellent product is obtained. A hide gains about 35 per cent in weight during the process of tanning. Other barks besides oak or hemlock are also used. So, for example, sumach is employed in making morocco leather from goatskins or imitation morocco from FIG. 71. A scene in a tannery. Scrubbing the hides. sheepskins. The tanning in this case proceeds rapidly; usually it lasts only a day or two. The hair of the goat- skins is removed by the lime process, but the excess of lime is then got rid of by using hen or pigeon manure. This keeps the leather soft. In other cases soft leathers are secured by removing the lime by means of sour bran prepared by treat- LEATHER, SILK, WOOL, COTTON, RUBBER 185 ing bran with water and sour dough. These methods of removing the lime are called " bating," for by employing infusions of the materials named the lime is removed and a large swelling of the hide is " abated." The bate liquor contains organic acids which form soluble salts with the lime, and these then can be washed out. Morocco leather is dyed with aniline dyes. Russia leather is produced in a similar manner, only the tanning is done with willow or tamarack bark. Kid leather, such as is used for gloves, is made of goatskins or sheepskins. These are unhaired with lime, and then treated with sour bran to remove the excess of the lime. The actual tanning consists of treating the skins with a mixture of common salt, alum, flour, and egg yolk in water. This is the so-called tawing process. The alum unites with the fibers of the skin. It is moreover antiseptic and so guards against putrefaction. The result obtained is a white, very soft, pliable leather. When the process of tanning is carried so far that the product no longer yields gelatine on being boiled in water, we have sole leather. The ordinary soft leathers yield some gelatine on boiling in water, but kid leather and chamois skin or buckskin do not. The latter are made from skins of deer, goats, or sheep in much the same way as morocco leather, only they are finally treated with animal oils to make them very soft and pliable. The excess of oil is removed by means of alkalies. Split leather is made by splitting thick hides, such as cowhide, by means of machinery. Three or four or more thin layers are thus made from one skin. These are then tanned separately and the grain is pressed upon them by machines. This split leather is cheap, but not as durable as that from natural thin skins. Much leather is tanned in America by the chrome-tanning process. The hides are unhaired and cleaned in the usual 186 CHEMISTRY AND DAILY LIFE manner. They are then soaked in a dilute solution of hy- drochloric acid and potassium bichromate, after which they are placed in a vat containing a reducing solution like sodium sulphite, for example. Thus hydrated chromic oxide is pre- cipitated, which unites with the fibers of the hide, much the same way that alumina does in making leather for kid gloves. The chrome tanning is completed in the course of a few hours. The leather is finally washed, dyed, and oiled as leather made with tannin. Box calf leather is chrome-tanned. Its cut edges are dark green in color, due to the presence of chromic oxide. Skins that are to be used for furs are tanned by the alum process. They are cleansed, dried, oiled, treated with sour bran and water, and finally tawed with a solution of common salt and alum. All alum-tanned, leather when thoroughly wet gets hard when it dries, for water dissolves out the alum, and the leather then dries hard like rawhide. Rubber is produced from the milky juice or so-called latex of the india-rubber trees, which grow in tropical climes. The best rubber comes from Brazil, and is called Para rubber. The milky juice is obtained from the trees by making inci- sions in the bark and collecting the exuding fluid in suitable vessels. About six ounces of latex are collected from a tree in the course of three days, and this yields about 1.8 ounces of rubber. A tree yields about ten pounds of rubber per year. The world produces about seventy-five thousand tons of rubber annually, of which supply about one-half comes from Central and South America, the tropics of Africa and Asia furnishing the other half in approximately equal shares. The island of Ceylon alone produces about 3000 tons per year. While the bulk of the rubber is still obtained from trees that grow wild in tropical forests, a fair share is already supplied by cultivated trees of rubber plantations. LEATHER, SILK, WOOL, COTTON, RUBBER 187 The method of preparing rubber from the latex is still rather primitive. After the juice has been collected, it is spread upon a piece of wood, about four feet long, shaped like a paddle. Usually the paddle is simply dipped into the liquid. It is then dried by holding the paddle over a fire so that the smoke comes into contact with the drying latex. Frequently the shells of oily tropical nuts are used as fuel, and the creosote and oily va- pors contained in the smoke of the fire perme- ate the material on the paddle and help to pre- serve it. When the first layer is dry, the paddle is again wet with the juice. This in turn is dried on over the fire, and so on, till many lay- ers have been formed and a thick covering has been obtained. This is then slitted and stripped from the paddle. It is called caoutchouc, and is placed on the market. This material is crude rubber. It is essentially a hydrocarbon corresponding to the composition (CsHs)*, and is probaby a condensation product of the hydrocarbon iso- prene, C 5 H 8 . Turpentine, Ci Hi 6 , is a related substance. FIG. 72. One method of tapping rubber trees. 188 CHEMISTRY AND DAILY LIFE Pure Para rubber consists of about 94.6 per cent caoutchouc, 0.14 per cent ash, 0.85 per cent water, 2.66 per cent resin, and 1.75 per cent protein substances. It is very elastic and not readily attacked by ordinary chemical agents. It does very gradually undergo oxidation on exposure to the air, and the oxygen thus taken up unites with the caoutchouc to form a hard, brittle compound. Thus it is that all rubber gradually deteriorates, whether in use or not. The rubber goods in the market are practically all made of so-called vulcanized rubber. The process of vulcanizing rubber was discovered by Goodyear in 1843. It consists essentially of adding about 10 per cent of sulphur to the caoutchouc, then molding the desired article from this mixture, and heating out of contact with the air to 140 to 150 C. A portion of the sulphur unites chemically with the caoutchouc ; the remainder, how- ever, is nevertheless required to produce the desired product. Such vulcanized rubber is more elastic, less porous, and not at all sticky as compared with caoutchouc. It is also less soluble in solvents like carbon bisulphide, turpentine, benzine, and naphtha, which are the common solvents for caoutchouc. Furthermore, vulcanized rubber does not lose its elasticity when subjected to cold. Sometimes rubber is vulcanized without heating, by simply treating it with chloride of sul- phur, which is a liquid at room temperatures. This is, how- ever, not as satisfactory for many purposes. White rubber articles are made by incorporating chalk or kaolin mechani- cally with the rubber. Red rubber is similarly colored by using antimony sulphide, black rubber by means of lamp- black, blue rubber by using ultramarine blue, and so on. Rubber coats, boots, shoes, etc., are made by spreading upon the cloth fabrics a plastic layer of caoutchouc properly mixed with sulphur. The vulcanizing is then accomplished by heating to 140 to 150 C. Treating with a solution of LEATHER, SILK, WOOL, COTTON, RUBBER 189 chloride of sulphur in carbon disulphide is also in vogue, but it does not yield as good results. Hard rubber, also called vulcanite or ebonite, is obtained when caoutchouc is mixed with about thirty to fifty per cent of sulphur and then heated to temperatures that are from thirty to forty degrees higher than tliose employed in mak- ing soft vulcanized rubber. Hard rubber is used in making combs, buttons, knife handles, bowls, trays, and many other useful articles. Rubber is a non-conductor of electricity and hence is much used as an insulating material for electrical wires. Hard rub- ber, too, is used as an insulator on electrical machinery and appliances. Gutta-percha, which comes from the gum of the percha tree, is chemically very similar to rubber. It is, how- ever, not as elastic. Being a good insulator, it is much used for covering wires and cables. It also serves like hard rubber for the manufacture of various articles. The demand for rubber has increased enormously of recent years on account of the use of rubber tires on automobiles and other vehicles. Attempts have consequently been made to make rubber artificially. While these have been success- ful, the synthetic rubber thus produced is nevertheless still so costly that it is not yet made on a commercial scale. Rubber substitutes are, however, on the market. These consist essentially of products obtained by heating certain oils, notably corn oil, to higher temperatures, where partial decomposition ensues and a somewhat elastic mass is ob- tained, which is worked up with sulphur. Such rubber sub- stitutes are not at all the equal of real rubber, and they are in general only employed when mixed with some of the gen- uine article. 190 CHEMISTRY AND DAILY LIFE QUESTIONS 1. What are the different sources from which the material for making wearing apparel is obtained ? 2. How distinguish between wool and cotton ? 3. What is shoddy ? 4. What is linen? 5. How are fabrics dyed ? 6. What is a mordant ? Describe its use. 7. Describe the process of making leather by use of oak tanbark ? 8. What is sole leather ? 9. What is split leather ? 10. How is leather for white kid gloves made ? 11. How is " buckskin" leather made? 12. How are skins that are to be used for furs treated ? 13. Describe the process of chrome tanning. 14. How is rubber prepared from the rubber trees ? 15. What is meant by vulcanizing rubber ? How does this process improve the material ? 16. How is hard rubber obtained ? What is it used for ? 17. What is gutta-percha, and what are its uses ? CHAPTER XIV THE SOIL LARGE areas of the earth's surface are covered with a loose layer of sand and clay mixed with the remains of plants that have decayed. This layer is called soil. Its depth varies greatly, being usually only from six to twelve inches, though soils may at times be several feet deep, as the experience that has been gathered incidentally in digging wells has shown. Living organisms of various kinds and variable amounts of water and air are also present in soils. The soil, which is the surface layer, rests upon the subsoil, which differs from the upper layer in that it contains less organic matter and is generally lighter in color. The subsoil gradually grades off into the solid rock below, as is illustrated by Fig. 73. FIG. 73. Soil formation showing the transition of rock to soil. 191 192 CHEMISTRY AND DAILY LIFE Sand and clay, which form by far the larger portion of soils, are mineral matter consisting of particles of silica, sili- cates, and frequently also carbonates, that have been broken off from large masses of solid rock. This breaking down or disintegration of solid rock into small fragments is very im- portant in the process of soil formation. It proceeds to- day just as it has been going on in the past. The chief agen- cies which are active in the process of making soil are water, wind, animals, and plants. The effect of each of these will now be considered. The work of water and the air in disintegrating rocks is com- monly termed weathering of rock. It proceeds constantly. The action of water and air is in part purely physical, that is, mechanical, and again it is partly chemical in its nature. All rocks are more or less porous and consequently allow water to soak into them. As this freezes, it expands and bursts the rocks, which then gradually crumble. Again, the water dissolves and carries away some of the material of which rocks are composed. This effect is chemical in character, and it further weakens the structure of the rock and also reduces the size of the particles. The waters of shallow streams roll the gravel stones along their beds. As these pebbles are carried along, they strike against one another ; their corners are knocked off ; they are ground smooth by the continual rubbing; and at the same time they are also reduced in size by the solvent action of the water. Thus the particles become smaller and smaller till they are nothing but sand. These finer particles then are generally the less soluble portions of the disintegrated rocks, and as they are carried forward by the stream they finally arrive in still water, where they sink to the bottom. When the stream eventually changes its course, they appear as real sand. In a similar way the waves on the shores of THE SOIL 193 lakes and seas are continually reducing rocks to fine particles. In the form of ice, water is a very important mechanical agent in the disintegration of rocks. So the cakes of floating ice as they are carried down by rapid streams in spring are com- monly covered with mud, rocks, and not infrequently por- tions of trees and other detritus that has been torn away from the banks. The ice on the shores of lakes incloses bowlders and shoves them upon the beach during the spring. Geologists have gathered evidence which shows that centuries ago the North American continent, about as far south as the Ohio River, was covered with ice to a great depth ; see map, Fig. 74. These immense glaciers moved southward like great rivers, only extremely slowly. Their movement was doubt- less entirely similar to that of the glaciers that exist at the present time in Switzerland, Greenland, and in the Rocky Mountains. This large field of glacial ice carried everything before it. It gathered up great rocks, which froze into the ice, and thus this heavy rock-shod surface slowly moved over the underlying rocks, crushing and grinding them to fine powder, which is now the sand and clay of the regions that have thus been visited by the glaciers. The latter have thus been invaluable in preparing the land for agricultural purposes. All over the North where the glaciers have passed may be found huge bowlders that escaped crushing and were left behind when the ice melted. These are of very hard material and so are commonly called " hard heads." They are round and smooth because of the grinding and rolling action to which they were subjected. It is clear, then, that the soils of these glaciated regions have really been transported to their present localities by the glaciers. In the South, on the other hand, where the glaciers have not been, the soils have not been transported. Here they have been formed from the 194 CHEMISTRY AND DAILY LIFE underlying rock upon which they rest to-day. Such soils that have been formed by the weathering of the particular rock upon which they rest are called sedentary soils, whereas FIG. 74. Map of North America, showing the southern limit of the glaciers. soils that have been carried from other localities to the place where they are found are termed transported soils. Thus the soils made by glaciers belong to the latter class. But transported soils are also found in regions that have not been glaciated, for wind and water are active everywhere in transporting soil material. Thus the rich soil in the lower THE SOIL 195 regions of river valleys consists largely of material which has been brought down by the river from the higher portions of the valley; and since this material has in many cases been brought from various rock formations, the resulting soil gen- erally possesses a greater fertility than if it had been formed by the weathering of any one kind of rock. The rich bottom lands of many of our rivers are excellent examples of this. A good overflow from a river is often as beneficial as a covering of manure. It should, however, be remembered that soil once formed is liable to be washed away. Soils composed of very fine material are thus carried off most readily, because the rain water does not soak into them rapidly, but accu- mulates on the surface and runs off in little streams which carry the fine topsoil with them. This is a decided loss for the farmer, for it is the best part of the soil that is thus washed away. In the formation of soils, the work of the wind consists largely in carrying and distributing particles of the soil ; but it nevertheless also takes part in the process of actually dis- integrating rocks and grinding coarser grains to finer ones. So when particles of sand are blown against solid rock, the latter is gradually worn away and becomes soil. The action of the wind in blowing particles of sand against rock surfaces is comparable with the effect produced by the water of a stream as it causes sand and pebbles to grind on the bottom of its bed. Sand dunes, like those along the shores of Lake Michigan, for example, are good illustrations of wind-borne materials. While these particular dunes are worthless for farming, there are nevertheless large areas of wind-formed soils, as, for example, the loess just west of the Mississippi River, which are quite fertile and deep. The air also acts chemically in the process of soil formation. Many rocks and rock particles that contain ferrous silicates are oxidized to 196 CHEMISTRY AND DAILY LIFE ferric compounds by the oxygen of the air, and thus farther disintegration results. This action is especially facilitated by water, which always contains oxygen in solution. Carbon dioxide, too, is always present in water that has been in contact FIG. 75. Sand dunes on the shore of Lake Michigan. with the air, and this dissolves many rocks, especially those that consist of the carbonates of calcium, magnesium, iron, etc. Animals are also important agents in soil formation. So> for example, rabbits, moles, prairie dogs, and other burrowing animals dig into the earth and throw out raw subsoil and pieces of rock, all of which help to make new soil through the further action of alternate freezing and thawing. Even the action of those humble creatures, the earthworms, is of material consequence in the work of forming soils. They bring a portion of the subsoil to the surface, draw dead leaves and other vegetable material into their burrows, and pass large quantities of the soil through their bodies, de- positing it on the surface at a rate estimated as ten tons per THE SOIL 197 acre annually. Ants are also active soil formers, and in some warm climates, as in Africa, they perform much the same work as the earthworm. Growing plants send their roots into rocks and soils, thus loosening them up and rendering them porous, a condition that allows air and water to enter more readily. Roots sometimes penetrate the soil to great depth, and as they de- cay after the death of the plant, they leave in the soil little channels through which water gains access to the lower strata. This water always contains carbon dioxide from the air ; and as already stated, it dissolves portions of the soil and thus makes them available as food for the rootlets of new plants. The roots and root hairs of living plants also give off carbon dioxide, which is absorbed by the soil water and then exer- cises similar solvent action upon the soil. Indeed, this car- bonated water in the soil is probably the chief agent that dissolves rocks and soil materials, thus enabling plants to obtain the necessary elements for their growth. It must not be for- gotten that all of these effects of plants in rock disintegra- tion and soil formation are slight in any one year; they never- theless amount to a great deal when carried on through hundreds of years. A small crack in a rock enables the tiny plants to work their way into it. As they grow, they exert great pres- sure, which is frequently sufficient to split the rock asunder. It is not uncommon to see bowlders weighing several tons split in two and a tree growing up between the pieces. Peat bogs are formed in swampy places or shallow ponds where plants grow year after year, dying at the end of the season, and falling to the ground or into the water. As this process is repeated for many years, there finally results a large accumulation of vegetable material. This rots more or less and so comes to be a type of soil. If the rotting takes place under water, it proceeds very slowly and results in the 198 CHEMISTRY AND DAILY LIFE formation of peat, which is fairly solid and generally shows some of the original shapes of the plants. As stated in Chap- ter IX, this material is in the first stage toward the formation of coal. In the low and undrained marshes of this country there are countless tons of peat waiting to be used for fuel or other purposes. Peat has been an important fuel in Europe for hundreds of years. If the decaying vegetable matter is exposed to the air and becomes alternately water-soaked and dry, it decomposes rather rapidly and forms muck. This is soft and spongy and does not show any trace of what it formerly was. It is usually mixed with sand and clay which have been washed over it. When a small quantity of soil is viewed with the naked eye, it appears made up of small particles which look like little grains. They are tiny pieces of disintegrated and de- cayed rock which range in size from those that may be read- ily seen to those that are as fine as dust. Besides these rock particles, which are the mineral matter, there are also present pieces of roots, stems, leaves, etc., which are the organic mat- ter. In peat soil the organic matter predominates. The rock particles that are large enough to be readily seen are called sand, while the very fine dustlike particles are termed silt and clay. The latter represents the very finest material. A magnifying glass is required to see the finest sand particles and to distinguish the silt from the clay. There are several important classes of soils that are commonly recognized. So when the sand particles are abundant, the soil is called a sandy soil. If the sand grains are not abundant, and the soil is quite floury when crushed, it is a clay soil. When there is a considerable proportion of sand present but also some clay, the soil is called a loam. The popular expressions sandy, light sandy, loam, heavy clay loam, and clay, all express the relative proportions of clay, silt, and sand in the soil. Loams THE SOIL 199 generally make good soils. They are usually fertile and easily worked. The vegetable matter in sands, loams, and clays is called humus, and muck soils are sometimes spoken of as humus soils. Real humus is vegetable matter so completely decayed that one cannot tell what it was originally. Though thus made up of minute rock particles and tiny pieces of decayed roots and stems, the soil nevertheless con- tains the substances necessary to make plants grow and develop to maturity. When a seed is placed in soil of proper moisture content, it absorbs water and germinates. Firming the soil about the seed brings it into closer contact with the moisture and hastens the absorption of water. But besides moisture, germinating seeds require oxygen, and so while the soil must be in close contact with the seed and contain sufficient moisture, it must not be water-logged nor so closely packed around the seed as to exclude air, otherwise the seed will fail to germinate and will actually rot. After germination has taken place a seed will continue to grow for some time at the expense of the reserve food which it contains ; but sooner or later growth will cease unless cer- tain chemical elements are available for absorption by the plant. These so-called essential elements without which the plant cannot grow are nitrogen, phosphorus, sulphur, potas- sium, calcium, magnesium, iron, and hydrogen and oxygen in the form of water. Plants indeed also contain other ele- ments like sodium, chlorine, and silicon, but these are gen- erally considered as non-essential, for plants have been grown without them. It should be clearly understood that the chemical elements mentioned are absorbed by the rootlets of the plants in the form of salts that are in solution in the soil water. Humus soils contain a great deal of carbon, but the plant gets its carbon content from the carbon dioxide of the air through its leaves (see Chapter IX), and not from the soil through its 200 CHEMISTRY AND DAILY LIFE roots. Some of the elements are used to build up the organic structure of the plant, while others remain in the plant as salts. So while hydrogen and oxygen unite to form water, hydrogen, oxygen, and carbon combine to form starch, cel- lulose, sugar, fats, and oils (see Chapter IX). Nitrogen, sulphur, and phosphorus together with carbon, hydrogen, and oxygen form the very complex substances called proteins (see Chapter IX), which are absolutely necessary for the growth of the plant. Each crop taken from the soil removes a certain amount of the essential elements which the soil contains. The amounts of such elements thus removed have been carefully ascer- tained by chemical analyses. These results, together with the chemical analyses of the soils themselves, have shown that most soils contain enough of the essential elements to furnish crops for hundreds of years. Nevertheless, as will be further explained below, the available supply of these elements is so limited that additional amounts must be added either as com- mercial fertilizers or as farm manures, if crop production is to continue at a profit. Chemical analyses have also shown that an acre surface foot of soil may contain enough potas- sium to last 1500 years, and supplies of phosphorus sufficient for 200 to 500 years, of nitrogen for 100 to 300 years, and sulphur for 200 to 300 years. The amounts of plant food removed from an acre by some of our common crops are given in Table 8. Besides the essential elements above mentioned, plants require large quantities of water throughout their period of growth. This water is necessary to dissolve the soluble salts in the soil and to convey them through the walls of the rootlets into the plant. Water is also needed to keep the cell walls of the leaf moist so that it may absorb carbon dioxide, and regulate the temperature of the plant ; for plants THE SOIL 201 JEAN BAPTISTS BOUSSINGAULT, 1802-1887. The first chemist who analyzed crops and manure and studied closely the relations of chemistry to farming. 202 CHEMISTRY AND DAILY LIFE TABLE 8 AMOUNT OF PLANT FOOD REMOVED FROM THE LAND BY VARIOUS CROPS, EXPRESSED IN POUNDS PER ACRE KIND OF CROP DRY WEIGHT OF CROP N p s K Ca Mg Ib. Ib. Ib. Ib. Ib. Ib. Ib. Wheat (grain), 30bu. 1530 34 6.2 2.5 7.7 0.7 2.1 Wheat (straw) . . 2653 16 3.0 3.7 16.2 5.8 2.1 Total crop . . 4183 50 9.2 6.2 23.9 6.5 4.2 Barley (grain), 40bu. 1747 35 7.0 2.6 8.0 0.8 2.4 Barley (straw) . . 2080 14 2.0 3.0 21.5 5.6 1.7 Total crop . . 3827 49 9.0 5.6 29.5 6.4 4.1 Corn (grain), 30 bu. 1500 28 4.3 2.5 5.4 0.3 2.0 Corn (stalks) . . 1877 15 3.5 2.2 24.7 7.5 3.3 Total crop . . 3377 43 7.8 4.7 30.1 7.8 5.3 Red clover hay . . 3762 98 10.8 6.1 69.2 63.9 16.9 Alfalfa hay . . . 9000 197 17.4 25.9 176.7 151.1 31.8 Turnips (roots) 3126 61 9.7 23.1 89.6 18.1 3.4 Sugar beet (roots) . 4320 53 8.8 3.8 72.9 7.1 7.8 Potatoes (tubers) . 3360 46 9.4 1.1 64.5 2.4 3.7 Tobacco (leaf) . . 1800 65 3.5 6.4 73.8 57.5 15.0 Tobacco (stalk) 3200 32 3.5 2.0 40.6 10.5 3.0 Total crop . . 5000 97 7.0 8.4 114.4 68.0 18.0 regulate their temperature by transpiration (i.e. exhalation of water vapor, from their leaves) just as animals keep their temperatures normal by means of perspiration from their skins. The amount of water thus required by a plant is very great. So, for example, it takes from 250 to 300 pounds of water to produce a pound of dry matter in the corn plant, and in the clover plant from 500 to 600 pounds of water are neces- sary to finally obtain a pound of dry material. The fertility of the soil depends upon a number of factors. THE SOIL 203 In order to be fertile a soil must first of all contain the chemical elements that are essential for plant growth. But these elements must also be available to the rootlets of the plants in the form of suitable soluble compounds. Now since nothing can pass into the roots except as it is dissolved in the soil water, it is clear that plant food must always be present in solution in the water of the soil in sufficient quantity in order that crops may thrive. Furthermore, undesirable bacteria and other injurious organ- isms must not be present, that is to say, the soil must be in proper sanitary condition. Again, a sufficient quantity of mois- ture, containing carbon dioxide to enhance its solvent power, must be present in the soil, and the latter must be sufficiently loose to enable the roots to penetrate it and air to enter and circulate in it. At the same time soil should not be so porous as to dry out too rapidly. Sandy soils commonly present the latter defect. As has already been stated, the solvent action of water upon the mineral matter of soils is rather slow, even when the water is charged with carbon dioxide ; and so, though a soil may contain a great abundance of the elements necessary for plant growth, this supply is often not avail- able in sufficient quantity as needed. The decaying or- ganic matter in the soil, to be sure, liberates carbon dioxide, which helps to increase the solvent power of the soil water, at the same time this rotted plant material furnishes to the soil in available soluble form the very elements that are essential for the growth of the crop. For these reasons the presence of considerable amounts of decomposing vegetable matter is essen- tial to the fertility of most soils. But often the entire supply of plant food in a soil is insufficient, and so it is necessary to fertilize the land, that is, to add material to it which contains abundant plant food in readily soluble, i.e. available, form. This is accomplished by treating the soil with manure or commercial artificial fertilizers that contain the ingredients 204 CHEMISTRY AND DAILY LIFE which are lacking. Thus from the decomposition of mineral particles, from the remains of roots and other vegetable matter left on the soil, as well as from manures and fertilizers used, there gradually accumulates a considerable supply of material that serves as food for the season's crop, which as it is taken from the land in turn leaves its roots and stems to decay, and so, aided by additional manure and fertilizer, furnishes the material necessary for the next crop, and so on. But there are always losses of plant food by erosion, the leaching of manure, and the sale of grain, sheep, cattle, hogs, milk, and other products from the farm. These constantly cause a re- duction of material for any succeeding year's crop. On the other hand, there may be very substantial gains made, especially of desirable compounds of nitrogen and carbon, by raising crops that absorb nitrogen from the atmosphere, a process which is accomplished by the action of bacteria growing on the roots. Clover, alfalfa, beans, peas, as well as other members of the legume family, are crops of this character. It is, of course, not always easy to estimate the extent of the losses which a soil has sustained, especially by erosion and waste ; nor is it simple to ascertain the gains that have been made by growing legume crops. Usually, the organic matter and essential elements present in available form, together with the soil reaction, are the most prominent factors that are to be considered in crop produc- tion. The humus produced by the decomposition of organic matter in the soil forms a coating around the soil grains. It helps to bind the soil grains together so that they are less readily blown about by winds. It also increases the moisture- retaining power of the soil, so that it can hold from two to three times its own weight of water. Again, in clay soils it lessens the tenacity with which the particles of the soil are held together, and this makes it easier to work these soils THE SOIL 205 properly. Humus is, moreover, a storehouse of nitrogen compounds, and consequently humus soils are always rich in nitrogen. By the slow action of bacteria, complex nitrog- enous compounds in humus are successively broken down into the simple compounds, ammonia, nitrites, and nitrates. The nitrates are the final and highest oxidation product. They serve as a source of nitrogen for the growing plants, whose roots absorb the nitrates dissolved in the soil waters. The process by which ammonia is transformed to nitrates in soils is called nitrification. The amount of nitrates in the soil waters at any one time is very small, yet it is very important, for most plants take their supply of nitrogen from the soil in this form. Plants utilize the nitrogen thus obtained in building up those wry important and complex constituents, the proteins. The preparation of the available supply of nitrogen for plants by the bacteria of the soil forms one of the most inter- esting chapters in the story of the feeding of crops. When plants die, their roots, stalks, and leaves are returned to the soil, where their decomposition is brought about by molds, bacteria, and other low forms of life. Much plant material is fed to animals, but even then a large residue of unoxidized matter is returned to the soil as manure. This vegetable matter or animal refuse is attacked by putrefactive bacteria, which reduce the proteins to simpler compounds. Still other bacteria convert these into ammonia, which really represents the first step in the process of nitrification, i.e. nitrate making. Ammonia is changed to nitrites by the nitrite bacteria, and finally the nitrate bacteria effect the oxidation of nitrites to nitrates. The bacteria do not really produce the nitrites and nitrates, but they oxidize ammonia to nitrous and nitric acid. As these acids accumulate, they stop further action on part of the bacteria ; hence the neces- 206 CHEMISTRY AND DAILY LIFE sity of having limestone or air-slaked lime present in the soil to neutralize the acids as they form. In this way calcium nitrate is left in the soil as the final product of nitrification. In the form of nitrates, the nitrogen then passes from the soil into the plant, where it is utilized in building up proteins. These are then returned to the soil direct, or they are first fed to animals, in which case the nitrogen is excreted in sim- pler compounds (urea, uric, and hippuric acids) in the urine and more complex ones in the feces. When returned to the soil, all of these compounds are first converted to ammonia and finally again to nitrates, and thus the nitrogen cycle completes itself over and over again as long as the sun furnishes the necessary energy for plant growth. While plants in general do not get their nitrogen supply direct from the air, it has already been mentioned that legumes are able to store up nitrogen from the air in little nodules on their roots, with the aid of certain bacteria. So, for example, the nodules on the roots of red clover, beans, and cowpeas are a matter of common observation. In these nodules are bacteria which take the nitrogen direct FIG. 76. Nodules on the roots from the air, and then give it of a soy bean. over to t h e p l an t i n the form of nitrogenous compounds. This process of converting the nitrogen of the air into nitrogenous compounds is commonly THE SOIL 207 spoken of as the fixation of nitrogen. The plant in turn gives the bacteria such food as sugar to live on. When the plant dies, there is left in these nodules on the roots a con- siderable supply of nitrogen, which upon subsequent nitrifi- cation becomes available for the next crop Thus it is that the growing of a good crop of cowpeas or soy beans will usually leave enough available nitrogen in the soil for a good crop of cotton or corn. Consequently raising a crop of a legume in a rotation is one of the most important methods of maintaining the available nitrogen supply of the soil. Soils are spoken of as acid, alkaline, or neutral according to their reaction toward litmus pamper. Acid soils are unsuitable for the successful growth of legumes, and it is therefore usu- ally desirable to correct such acidity by an application of fresh lime, ground limestone, or marl. It is better to use old, thoroughly air-slaked lime, rather than fresh quick- lime, for the latter has caustic properties. A ton of finely pulverized limestone per acre once in five years is commonly not too much. The amount of phosphorus in most soils is generally small, being from 0.02 to 0.1 per cent. All crops take phosphorus from the soil, in which this element is present as phosphates. Plants store up phosphorus, especially in their seeds. The supply of phosphates in the soil can consequently be main- tained by lessening the sale of seed, or crops, or by adding suitable phosphorus containing fertilizers. If a farmer buys plentifully of feeds, especially bran, for his stock and sells only milk and animals, he may not need to purchase phos- phorus in the form of commercial fertilizer. Indeed, the soils of farms of dairy states, where milk and cream are sold and feeds are purchased, may actually show a gain of phosphorus. Raw rock phosphate, also called " floats," is a cheap commer- cial fertilizer for soils deficient in phosphorus. When used 208 CHEMISTRY AND DAILY LIFE with organic material, although rather slow in becoming available on account of the fact that its phosphates do not dissolve very readily, it nevertheless gives good results. Superphosphate, or acid phosphate (which, as has been stated in Chapter VII, is made by treating bones or phosphate rock with sulphuric acid), dissolves far more readily and hence is immediately available to plants. It should be used when a supply of organic matter is not at hand. There is usually a sufficient amount of available potassium in soils. This is especially true of clay soils, which are rich in feldspars. And so when these soils are furnished with plenty of organic matter that gives off carbon dioxide as it decomposes, thereby increasing the solvent action of the soil water on the feldspathic constituents, an ample supply of potassium is furnished to the rootlets of the growing crop. Some soils, like sands and marshes, are deficient in potassium. To these it should be added at the rate of fifty to one hundred pounds of potassium chloride or potassium sulphate per acre. Although the quantity of sulphur in soils is usually low, being about as low as that of phosphorus, it is replenished to a considerable extent during the year by rains. These carry sulphur dioxide down in solution from the atmosphere. The sulphur dioxide has come into the air mainly from the burning of coal. For some crops which need much sulphur, like onions, cabbage, turnips, and rutabagas, additional sulphur in the form of gypsum, calcium sulphate, is advantageous. From 100 to 200 pounds per acre is a good application. It has already been stated that the soil must be in proper mechanical condition so that it will readily hold the proper amount of water, permit air to circulate in it, and allow the roots to penetrate it. Proper tilth of the soil, that is to say, proper mechanical working of the soil with the plow, spade, THE SOIL 209 hoe, harrow, roller, etc., is of great importance. While it is desirable that the particles of the soil be sufficiently fine, it must also be remembered that they may be so fine and closely compacted together that neither air nor root hairs can gain entrance. This condition is just as unfavorable as a coarse, lumpy soil. In either case the water storage capacity is decreased. From soils that are too loose and porous, moisture drains away, and the air circulates so freely in them that they dry out too rapidly. It has already been stated that this is a common defect of sandy soils. By the force of capillarity, or capillary attraction, water rises in tubes of small bore or between particles of solid substances. Fine-grained soils, of which clay soils are an excellent example, have small pores or spaces between their particles, and hence water will rise from below and be nearer to the surface in these soils than in those of coarser grain. It is indeed possible to aid the upward capil- lary movement of water so that in a dry season, when the seeds have been planted, moisture will be drawn to them from be- low. This is accomplished by rolling the ground, for roll- ing forces the particles of soil closer together and thus a stronger upward movement of water by capillarity results. Much water that has thus been drawn up by capillarity is lost by evaporation at the surface of the soil. This loss can be greatly reduced by covering the surface with a protecting layer or soil mulch. So if the surface of the soil to the depth of two or three inches is thoroughly stirred on a dry, windy day, the dry layer that is thus formed becomes an effective barrier against large losses of moisture from below. In regions where rainfall is light this practice is of the very greatest importance. It is upon this principle that so-called dry fanning is based. In speaking of soil waters, the terms gravitational water, capillary water, and hydroscopic water are often used. Their 210 CHEMISTRY AND DAILY LIFE meaning will be clear from the following experiment. If the neck of a funnel be stoppered, and the funnel be filled with soil, and then water poured on till the funnel is brim full, the soil will be in a saturated condition. If now the stopper is- removed, a considerable portion of the water will drain off ; this is called gravitational water, for it runs off because of the action of the force of gravitation. If the lower end of the funnel is now stoppered again and the whole allowed to stand, with the upper surface exposed, a large part of the moisture retained by the soil will rise up to replace that which is lost at the surface by gradual evaporation. This upward movement of the water in the soil is due to capillarity, and the water held and moved in this way is termed capil- lary water. If the funnel is thus allowed to remain exposed to the air for several days or weeks, the soil it contains will become dry. But if a portion of this apparently dry soil be weighed out and then heated at the temperature of a boiling water bath, 100 C., it loses still more moisture. The water thus expelled is called hydroscopic water. It is obvious that it is held more tightly by the soil particles than the capillary water. The removal of gravitational water by proper natural or artificial underdraining is necessary because, if it remains in the soil, the latter becomes water-logged, and so displaces the air which is necessary in the soil for the growing roots. Drainage lowers the water level in the soil, causing roots to penetrate deeper for moisture, and consequently their feeding ground is increased. This helps them to withstand drought. The amount of moisture in a soil must be taken into consider- ation in determining whether it is an opportune time to work the soil or not. So, for instance, if heavy clay soils are worked when wet, the particles are forced close together and " pud- dling " is the result. This is a brickmaker's term. In making THE SOIL 211 brick the first effort is to destroy the granular texture of the clay, which is accomplished by wetting and working it. If this is done with clay soil, it obviously will be left very com- pact and hard when it dries. Such a condition, once brought FIG. 77. Cracked soil. Large losses of water will result from such a condition. about by working the clay soil when too wet, may last for a number of years, and so it is wise for the farmer not to get his clay fields puddled. When land breaks up cloddy, it is desirable that it be plowed long before planting in order that freezing, thawing, and weathering may break up the clods and make them loose and mellow. This requires time, but for best results long exposure to the weather will pay. The 212 CHEMISTRY AND DAILY LIFE effect of humus on fine-grained soils, such as clays, is to divide the particles and so lessen the tendency to puddle. In sandy soils, humus tends to cement the grains of sand together more firmly, a result that could not be secured by mere films of water around such coarse particles. The temperature of soils is affected by their power to ab- sorb heat from the rays of the sun. The readiness with which heat is thus absorbed is greatly influenced by the color of the soil. Thus dark-colored soils absorb heat readily and so warm up relatively rapidly. They are consequently called warm soils. On the other hand, light-colored soils do not absorb heat so readily as dark-colored ones of equal moisture content, and they consequently do not warm so quickly. They are called cold soils. It must be borne in mind, however, that the amount of moisture in the soil is the most important factor in determining the rate at which it will warm up, for it takes large amounts of heat to evaporate water, and so dry soils always warm up more rapidly than wet ones, regardless of their color or composition. Barefooted boys well know that dry sand and fine road dust become warm more quickly than a wet soil. Indeed, the evaporation of water from the soil cools the latter, just as the evaporation of sweat cools the animal body. Dark-colored, sandy soils warm up earliest in spring because the gravitational water soon drains out of them and they readily absorb heat from the sun's rays. Such soils, then, are most suitable for market gardening or early spring crops. THE SOIL 213 QUESTIONS 1. From what is soil made, and what are the agencies active in its formation ? What is a glacier ? 2. How are peat bogs formed ? 3. What is the difference between sedentary and transported soils ? 4. What is the difference between a sandy soil and a clay soil ? What is humus ? 5. Name the essential elements needed for crop production. 6. About how much phosphorus is there in a soil, and how many crops of wheat would it produce ? 7. How much water is required to make a pound of dry matter in a corn plant ? 8. Why is it necessary to have plenty of organic matter in the soil? 9. How would you maintain the nitrogen content of soils ? 10. What may the reaction of the soil be, and what should be the proper reaction for the growth of legumes ? 11. What is meant by the puddling of clay? What is the in- fluence of humus on the water-holding power of a soil ? Distinguish between capillary and gravitational water. 12. How does drainage help a plant in a season of drought ? CHAPTER XV COMMERCIAL FERTILIZERS COMMERCIAL fertilizers are manufactured plant foods. They are usually kept for sale at warehouses and seed stores. About $100,000,000 is spent annually in the purchase of fertilizers in the United States and probably one-half of this money is thrown away. This is not an argument against the use of commercial fertilizers, but it does mean that they should be purchased with proper judgment, and not used at all until trials have shown that they are actually necessary. Experimental investigations and common experience have shown that material increases in the yield of the land have in most cases been obtained by adding to the soil but three of the essential elements necessary for plant growth ; 'namely, nitrogen, phosphorus, and potassium. These elements, it is to be recalled, are not applied in the free or uncombined state, but in the form of compounds, usually salts, that will dissolve in the soil waters, so that they can be absorbed by the rootlets of the growing crop. In consequence of the results of experience, then, commercial fertilizers as placed on the market to-day contain as their essential ingredients only nitrogen, phosphorus, and potassium in the form of mixtures of salts or other compounds of these elements ; and thus it is that the nitrogen, phosphorus, and potassium in a commercial fertilizer are the only materials which give that fertilizer value. It should be remembered, however, that equally good results are sometimes secured by applying to the land dressings that 214 COMMERCIAL FERTILIZERS 215 FIG. 78. Effect of fertilizers on a crop of potatoes. (^4) Complete fertilizer used. (B) No fertilizer used. 216 CHEMISTRY AND DAILY LIFE do not contain any of the three elements just mentioned. So, for example, lime or limestone and gypsum are fre- quently beneficial. The benefit derived from lime is usually due to its basic properties, which cause a neutral or alkaline reaction of the soil. Gypsum, on the other hand, acts as a source of sulphur, and also produces physical changes in the soil. Commercial fertilizers are made from a few basal materials which are articles of commerce. A so-called complete ferti- lizer consists of two or more such basal materials mixed to- gether so as to give the desired percentage content of nitro- gen, phosphorus, and potassium. These basal materials are derived (1) from inorganic or mineral matter, and (2) from organic matter, that is to say, from animal or vegetable material. So, for example, as sources of nitrogen, the inorganic salts, ammonium sulphate and sodium nitrate, and organic mate- rials like dried blood, meat scraps, tankage, fish scrap, and cotton-seed meal are in common use. As sources of phos- phorus, phosphate rock and basic slag, which are inorganic, and bones and guano, which are organic, are utilized. Fi- nally, as sources of potassium, potassium chloride, potassium sulphate, and kainite, all of which come from the Stassfurt salt deposits, are employed. All of these materials will now be considered more in detail. Among the nitrogenous fertilizers ammonium sulphate, (NH 4 ) 2 SO 4 , also popularly called " sulphate of ammonia," is used to a considerable extent. This salt has already been mentioned in Chapter IV. It is a by-product of the manu- facture of illuminating gas from coal, also of the dry distilla- tion of bones in the manufacture of bone black. Ammonium sulphate is a very concentrated fertilizer, containing 20 per cent of nitrogen. It is copiously soluble in water, does not readily leach out of the soil, and very quickly undergoes COMMERCIAL FERTILIZERS 217 nitrification, that is, conversion into nitrates. Sodium ni- trate, NaNO 3 , also called nitrate of soda, has been described in Chapter X. It occurs as Chili saltpeter in extensive de- posits in the rainless districts of western South America. The natural material from the saltpeter beds contains a large amount of common salt. This, however, is practically removed from the product before it is placed on the market, so that as sold it is a crude nitrate of soda containing from 15 to 16 per cent of nitrogen. Nitrate of soda is readily soluble in water and easily leached from the soil. It should conse- quently be applied just before planting or after the plant has possession of the soil. Fifty to one hundred pounds per acre is a proper application. Dried blood, meat scraps, and tank- age come from the slaughterhouses. Dried blood is simply blood of the slaughtered animals dried and ground. Like all organic nitrogenous fertilizers, it is insoluble in water ; but it ferments readily in the soil and yields its nitrogen in soluble form to the growing plant. Dried blood contains from 13 to 14 per cent of nitrogen, while meat scraps contain variable amounts of nitrogen, usually from 4 to 10 per cent. The value of fish as a fertilizer was known to our American Indi- ans. Squanto, an Indian, taught the New England settlers that they could increase the yield of corn by putting a fish under each hill. This gave the grain the two elements it needed most, phosphorus and nitrogen. To-day fish ferti- lizers are made in large amounts from rough fish and fish oft'al on the Atlantic coast and the Great Lakes. Cotton- seed meal is obtained by removing the hulls and oil from cotton seed. The residue is then ground and put on the market. It is extensively used as a fertilizer in the south. There are still other nitrogenous fertilizers like waste leather, hoof and horn meal, hair from the slaughterhouses, and wool waste from the woolen mills. The chemical nature 218 CHEMISTRY AND DAILY LIFE of these 'has already been described in Chapters IX and XIII. All of these materials decompose so slowly in the soil that many states have passed laws prohibiting their sale as fertilizers. In ordinary farming it is seldom profit- able to purchase nitrogenous fertilizers, for the nitrogen supply of the soil can be maintained by means of farm manure and the proper use of legume crops in rotation, as already stated in Chapter XIV. Such crops as the clovers, peas, and beans are aids to the farmer as nitrogen gatherers, especially if he occasionally plows under a goodly part of the plant. In intensive farming, like market gardening, it will be necessary to make liberal use of nitrogenous fertilizers of commercial origin. Since three-fourths of the phosphorus absorbed from the soil is deposited in the grain of the crop and therefore ordi- narily sold from the farm, it is clear that phosphatic ferti- lizers are of fundamental importance. In the soil, phosphorus occurs as the phosphates of calcium, magnesium, and iron. As already pointed out in Chapter VII, phosphates are salts of phosphoric acid, H 3 PO 4 . In all the important commercial phosphate fertilizers, including phosphate rock, bones, basic slag, and guano, phosphorus is present as phosphates formed by the combination of phosphoric acid, H 3 PO 4 , with lime, CaO. These are phosphates of calcium, commonly spoken of as phosphates of lime. With lime, phosphoric acid forms three important compounds : (1) insoluble phosphate of lime, which is tricalcium phosphate, Ca 3 (PO 4 ) 2 ; (2) soluble phosphate of lime, which is monocalcium phosphate, CaH 4 (PO 4 )2 ; and (3) reverted phosphate of lime, which is dicalcium phosphate, Ca 2 H 2 (PO 4 ) 2 . The first of these, as its name indicates, is almost insoluble in water and hence is not readily available to plants. Ground bone, guano, and floats contain their phosphorus in this form. Ground bone, or bone meal, is a product of the packing houses, COMMERCIAL FERTILIZERS 219 glue factories, and soap works, the raw material being the bones of farm animals. The bones are either ground directly or after they have been steamed and dried. In the latter case they are called steamed bone. Raw bone con- tains 2.5 per cent of nitrogen and 11 per cent of phosphorus. It is more effective when finely ground than when coarse ; furthermore, raw bone decomposes more slowly in the soil than steamed bone, because the fat has been removed from the latter. Guano, a highly prized but comparatively rare fertilizer, consists of the excrements and remains of sea fowls, which have accumulated in certain localities along the west coast of South America. It contains nitrogen, potassium, and phosphorus, the latter amounting to from 4 to 8 per cent. The phosphorus in guano has come from the skeletons of of the birds. Floats is the term applied to the finely ground FIG. 79. Phosphate-rock mining in Florida. crude phosphate rock found in Tennessee, the Carolinas, Florida, Georgia, and other Southern states. It has recently also been discovered in considerable beds in Utah and 220 CHEMISTRY AND DAILY LIFE Wyoming. It contains from 11 to 13 per cent of phosphorus and is the chief source and also the cheapest source of phos- phorus supply on the market. Recent investigations have shown that when used with sufficient organic matter, such as farm manure or green manure, ground phosphate rock has high fertilizing power, for the organic matter on decom- posing furnishes carbonic acid, which aids in dissolving the tricalcium phosphate, thus making it available. Soluble phosphate of lime is also known as " one lime phosphate," acid phosphate, acidulated rock, superphosphate, etc. It is made by treating bones or phosphate rock with sulphuric acid, as stated in Chapter VII, where the chemical equation explaining its formation is given. Because of its solubility in water, this fertilizer contains the phosphorus in the most available form for direct use by the plant. A good sample con- tains about 7 per cent of phosphorus, or only about half as much as a good sample of rock phosphate. Furthermore, a ton of the material contains about half a ton of gypsum. The latter is formed simultaneously when the sulphuric acid acts on tricalcium phosphate. Though readily soluble in water superphosphate is not leached out of the soil, for the lime an Q^ boric acid, sail- ofS o cyhc acid', for- ^^fl maldehyde, and .YJ^o- &.* .&J*^** o benzoic acid. .* *8& foS&Shft quantity is gen- 6 ~ ' u ^ erally prohibited FIG. lOO. Pure and impure milk. The long black bylaw; and this is right, because when used they show that uncleanly methods of milk production are being practiced. Milk produced under clean, sanitary conditions, and well cooled, will keep as long as it is necessary for transportation to the city and the con- sumer. The fat of milk exists in globules which tend to rise to the surface when the milk stands. After the fat has been removed, what is left is called skimmed milk. Cream is milk with a large amount of fat in it. Cream can be separated from milk by gravitation, or by the much greater centrifugal force produced by rapid rotation in the centrifugal separator. There are two ways of separating by gravity, namely, the shallow-setting and the deep-setting methods. In the shallow- setting, the milk is placed in pans two to four inches deep, cooled to 60 F., and allowed to stand from 24 to 36 hours. The cream is then removed with a slightly concaved ladle. Not all the fat is obtained in this way; indeed, only about 80 per cent of the fat is usually thus removed. In the deep- setting process, the milk is put into cans twenty inches deep 298 CHEMISTRY AND DAILY LIFE and less than one foot in diameter, which are then placed in ice-cold water. The cream rises rapidly and the operation is practically complete in 12 hours. In this way 90 to 95 per cent of the fat can be removed, depending upon the conditions of cooling, ma- nipulation, and the breed of the cow whose milk is thus treated. Because of its small globules, the fat of the Hoi- stein milk cannot be as completely removed as that of the Jersey milk by this process. Cream separa- tors are machines which are now in general use. In them the milk runs into a bowl which is revolv- ing several FIG. 101. -A modern cream separator. thousand times a minute. This revolving tends to throw the heavy par- ticles to the outside of the bowl. Since the fat is not as heavy as the other constituents of the milk, it tends to come to the center of the bowl. The cover of the bowl is so con- structed that the cream, or fat, can escape from the center of the cover, while the skim milk escapes from the opening near the edge of the bowl. There are many different makes MILK AND ITS PRODUCTS 299 of separators, but the principle of operation in all of them is practically the same as that just described. These machines recover from 97 to 98 per cent of the fat in the milk. A well-operated separator rarely leaves as much as 0.1 per cent of fat in the skimmed milk. A good cream usually contains from 25 to 30 per cent of fat. Skimmed milk varies in composition according to the more or less complete removal of the fat. It differs from whole milk only in its smaller content of fat, and so it still contains a val- uable amount of foodstuffs which should be used on the farm. It is excellent for feeding pigs. Though poor in fat, machine- separated milk has the advantage of being sweet and of keep- ing better than the product from other processes of skimming. Upon agitating cream or milk for some time, the fat globules coalesce and butter separates out in irregular masses. Churning is a mechanical process in which the fat globules of the cream collide and adhere; the large, irregular masses thus formed then become centers of growth to which still other fat globules adhere. To be of good quality, butter must possess a certain texture and grain and be neither hard nor smeary. This desirable result can be secured only by churning at a favorable temperature. If the temperature is too low, the butter will be long in coming and very hard. If the temperature is too high, the butter will come quickly, but it will be greasy and destitute of grain. The temperature of churning should not vary more than from 45 to 65 F. ; indeed, in most cases from 50 to 60 F. is chosen as the proper range. In the manufacture of butter, the cream is either allowed to sour (ripen) before it is churned or is churned directly without ripening. The former method yields the ordinary market butter, the latter method the sweet cream butter. In the ripening, acids and other products are formed, which give the butter a high flavor. Butter made 300 CHEMISTRY AND DAILY LIFE from sweet cream is usually not salted. It contains some- what more water, fat, and casein than sour-cream butter. Ordinary market butter contains about 12 per cent of water, 84.2 per cent of fat, 1.3 per cent of curd, and 2.5 per cent of salt. Under the Federal Pure Food Law, butter must con- tain 82.5 per cent of fat and not more than 16 per cent of water. During the working of the butter a part of the buttermilk is squeezed out. This buttermilk does not vary much in composition from the skimmed milk. Oleomargarine, which is much used as a substitute for butter, is made by churning together " oleo oil," " neutral oil," milk, and a small amount of cottonseed oil and peanut oil. " Oleo oil " is the liquid oil pressed out of beef fat. " Neutral oil " is melted lard. Some of the oleomargarine is colored yellow so as to imitate butter, in which case it is taxed 10 cents per pound. Renovated butter is made from old and rancid butter, by melting it, separating the fat from the casein, and blowing air through the fat to remove the un- pleasant odors which are due to volatile acids that have formed. The liquid fat is then churned with milk, and a granular mass is obtained upon cooling. This is then worked, salted, and made up as butter. Numerous kinds of cheese are made from milk. All of the casein, nearly all of the fat, and also a large part of the ash of the milk are present in cheese. The albumen and milk sugar are in the whey. Only two types of cheese which are com- monly made in this country, namely, Cheddar cheese and cottage cheese, will be described here. Cheddar cheese is made by adding a small amount of rennet extract to the milk. The rennet is made by extracting the fourth stomach of the calf with a dilute salt solution. Rennet contains a ferment which causes the milk to curdle. In the process of Cheddar cheese making, the milk is warmed to 84 F. and " ripened " MILK AND ITS PRODUCTS 301 to 0.25 per cent acidity. Rennet is then added, and when the curd is firm, it is cut into small cubes. In this process the fat has become entangled mechanically in the curd. The vat is now warmed, which causes the curd to shrink and harden. After being maintained warm for one to two hours, the whey is drawn off and the curd is piled. In this condition it mats into a solid mass. This is now passed through a grinding mill, FIG. 102. A cheese curing room. salted, and pressed into molds. The cheese is next put into curing rooms at a temperature of from 50 to 60 F. and al- lowed to ripen. When three to four months old, it is put on the market. The practice of ripening the cheese at a tempera- ture as low as 30 F. is coming into general use. The process of making cottage cheese is substantially as follows. Instead of adding rennet to curdle the milk, it may be allowed to stand until it is sour; or it may be soured by adding a " starter." A starter is a powder, containing living bacteria, 302 CHEMISTRY AND DAILY LIFE which will slowly produce lactic acid when added to milk. The curd thus obtained is strained from the whey, and after being salted, this makes sour-milk or cottage cheese. The following data will give a general idea of the composi- tion of cheese. Cured Cheddar cheese contains about 34 per cent of water, 35 per cent fat, 28 per cent casein and 3 per cent salt. Swiss cheese, which is made extensively in some parts of this country, has about the same composition. Some types of cheese, as Camembert and Roquefort, contain as high as 50 per cent of water. Milk is sometimes partly skimmed before being coagulated with rennet, and the prod- uct thus obtained is called skimmed-milk cheese. In many states its manufacture is prohibited. There are still other dairy products on the market. Con- densed milk is manufactured from whole milk or partly skimmed milk by evaporating off a certain portion of the water. The milk is evaporated in large vacuum pans to one- third or more of its original volume. Condensed milk should contain at least 8 per cent of fat. In many cases cane sugar has been added in large quantities. This aids in preserving the product, even after the cans are opened. A sweetened con- densed milk may contain 40 per cent of cane sugar. Con- densed milk is a good substitute for whole milk or cream, and it is used when fresh milk cannot be obtained, as on sea voyages, in mining camps, etc. Milk powders are made from skimmed or partly skimmed milk by allowing the milk to evaporate in a thin layer on a revolving drum. The product is scraped off in flakes, and placed on the market as thin, yellow scales. By a later process this powder is produced as a fine flour and this is destined to be a very efficient way of handling milk for long shipment's. Most milk powders cannot be made from whole milk, as the high fat content of the powder impairs their keeping quality. Ice cream is made from rich milk or from MILK AND ITS PRODUCTS 303 cream ; the latter is the better. Ice cream consists of milk or cream, sugar, eggs, flavoring material, and sometimes cornstarch, gelatin, or gum tragacanth. The last three sub- stances are added to give the ice cream smoothness and body. No two things have done more to revolutionize the dairy industry than the cream separator and a simple test for fat. FIG. 103. A Babcock tester run by steam. The cream separator is a European invention, but the simple efficient fat test for milk was invented in America by Pro- fessor S. M. Babcock of the Wisconsin Experiment Station in 1890. The Babcock test is a simple means of testing the milk for its amount of fat. By its use the dairyman can determine which of his. cows are paying their board, and in this way he can definitely ascertain the value of a cow for butter produc- tion. Knowing the per cent of butter fat in the milk and the quantity of milk produced, it becomes an easy matter to compute just how much butter fat is being produced. Again, 304 CHEMISTRY AND DAILY LIFE on the basis of the test, milk can be paid for on the basis of its fat content at creameries and cheese factories. Since the Babcock tester has become an established part of practically every dairy equipment, the old and but too common practice of watering milk has largely disappeared. Coupled with the Babcock test for the detection of water adulteration there should be used another equally important instrument, called the lactometer. This instrument resembles a floating ther- mometer (see Fig. 107). Its purpose is to determine the spe- cific gravity of milk. Normal milk has a specific gravity of 1.029-1.034, and if the fat is partly removed, the specific gravity is raised ; but if the adulterator now adds water, he may bring the specific gravity back to that of normal milk. The use of both tests detects such tampering with the milk. For a more just distribution of dividends at cheese factories, there is coming into use a simple mechanical way of estimating the amount of casein in milk. This is known as the Hart casein test; it was devised at the Wisconsin Experiment Station in 1907. This method is almost as simple as the Babcock test for fat, and it will help to solve the problem of proper payment for milk at cheese factories. With both the fat and casein tests the cheese-yielding capacity of the milk can be determined. QUESTIONS 1. What does the milk of animals come from ? 2. About how much fat is there in average milk, and how does it exist in the milk ? What is this fat used for ? 3. Which has the larger globules, Jersey or Holstein milk ? 4. What is the principal use of casein in cows' milk ? 5. How would you obtain the milk sugar from milk ? 6. What precautions should be taken to produce clean milk? What is meant by pasteurization, and sterilization ? MILK AND ITS PRODUCTS 305 STEPHEN MOULTON BABCOCK. 1843-. Scientist and inventor of the fat test for milk, which bears his name, test has revolutionized dairying. X This 306 CHEMISTRY AND DAILY LIFE 7. What is an antiseptic? Name two antiseptics and state whether they should ever be allowed in milk. Give reason for your answer. 8. Name three ways of separating cream from milk. How much fat is there in good cream ? 9. Why does butter sometimes fail to come ? How much water is there in ordinary butter ? 10. Describe how oleomargarine and renovated butter are made. 11. How is Cheddar cheese made, and what is its general com- position ? 12. Who invented the commonly used fat test for milk, and where was it done ? Of what importance is it ? What two tests should be used at cheese factories for payment of milk ? 13. What is a lactometer, and what is it used for ? CHAFPER XXI POISONS FOR FARM AND ORCHARD PESTS IN order to have vigorous, healthy plants and sound fruits, the insect pests that are ever present to feast upon them and so impair their growth or even destroy them must be exter- minated. To kill off such harmful insects that infest trees FIG. 104. Potato bugs and their eggs. and other plants, various poisons have come into use. These poisons are quite numerous, but they may all be divided into two great classes according to the type of insect that is to be destroyed. Thus there are pests, like the potato bug, that feed directly on the whole leaf and will consequently be killed by any poison on the leaf. Such poisons which destroy because they are thus eaten by the insects are called stomachic poisons. 307 308 CHEMISTRY AND DAILY LIFE Again, there are other insects like the aphides or plant lice and the San Jose scale that get their nourishment not by eating the whole leaf, but by merely sucking out its juices. To exterminate pests of this kind requires a poison that acts FIG. 105 (A). San Jose scale on an apple twig, slightly enlarged. FIG. 105 (B). San Jose scale on an apple twig, greatly enlarged. by enveloping the insect and thus corroding it, poisoning it by absorption through its skin, or cutting off its opportunity to breathe. Poisons that act in this manner are called con- tact poisons. Stomachic poisons generally contain arsenic, which is present not as the free element, but either as the oxide As 2 O 3 , POISONS FOR FARM AND ORCHARD PESTS 309 also known as white arsenic or arsenious acid, or as some more complex arsenical compound, like Paris Green or lead arsenate. The exact 'composition of these compounds has already been given in Chapter VII. White arsenic was first used as an insecticide ; but though not copiously soluble in water, it nevertheless dissolves sufficiently to yield a solution that corrodes or " burns " the foliage. It was therefore soon discarded. In fact, any arsenical poison that is to be used as an insecticide must have wry little uncombined white arsenic, As2Os, in it. Paris green is at ^present the standard poison for all insects that bite and swallow their food. Its composition is expressed by the formula Cu 3 As 2 O 6 * Cu(CaHsOi)2. This poison is prepared by adding a hot solution of arsenious acid to a hot solution of copper acetate. From this mixture Paris green separates out as a fine powder of a bright green color. Pure Paris green is almost insoluble in water, but it will readily dissolve in ammonia water, the resulting solution being dark blue. This is a test by means of which certain impurities in Paris green may be detected. Thus, if a sample of the sub- stance is adulterated with gypsum, as sometimes occurs, the latter will form a white suspension in the ammonia water and finally settle out on the bottom of the glass. If no solids separate out, it of course only shows that impurities which are insoluble in ammonia are not present. The glass test may also be applied. In this test a small amount of the Paris green is placed in a glass ; this is then inclined and gently tapped so that the poison will slowly move down the inclined plane. In the case of pure " green," the dust will be of a bright green color. If it is impure, it may yield a white, pale-green streak, depending upon the color of the adulterating substance that is present. Examination under the microscope will also help to determine its purity. Paris green contains about 58 per 310 CHEMISTRY AND DAILY LIFE cent of arsenious acid, 31 per cent of copper oxide and 10 per cent of acetic acid. These are all chemically combined in one compound, whose composition is indicated by the for- mula given above. Paris green should not contain over 5 per cent of free arsenious acid, or it will seriously burn the foliage. For use in spraying, from 6 to 8 ounces of Paris green should be thoroughly worked into a paste with a little water, and then added, to 50 gallons of water. In addition, about 2 pounds of lime should be added; this will neutralize any free arsenious acid present, and also act as a " marker " on the sprayed plants. London purple is an arsenical poison which was first im- ported from England as a substitute for Paris green. It is prepared by boiling with slaked lime a purple, arsenious acid bearing residue from the dye industry. Arsenite of lime is thus formed, together with some arsenate, which contains more oxygen than the arsenite. London purple is more injurious to foliage than good Paris green, because of its content of free arsenious acid. It should always be used with lime. Lead arsenate was recommended as an insecticide in 1892 and was first used against the tent caterpillar. It is pre- pared by adding lead acetate solution to sodium arsenate also dissolved in water. These substances dissolve readily in the cold and react to form sodium acetate and lead arsenate Pb 3 (AsO 4 ) 2 . This poison should be handled as a paste, for when once dried it does not again remain well in suspension. It is the most insoluble of the insecticides now in use. Fur- thermore, it adheres wry tenaciously to the leaves and is least liable to scorch them. For spraying purposes two pounds of the commerical paste in 50 gallons of water is the proportion com- monly used. Hellebore is often recommended as an insect poison. This is the ground root of the poke-root plant. Pyrethrum, or POISONS FOR FARM AND ORCHARD PESTS 311 insect powder, which comes from the flower heads of certain^ plants, is another poison of similar character. Both of these materials contain poisonous substances, but they deteriorate much with age. Contact poisons may act (1) by their caustic properties, (2) by poisoning because they are absorbed by the surface of the insect, or (3) by closing up the insect's breathing tubes. Lime-sulphur and kerosene emulsion are types of these poisons. They are commonly used against the scale insects. Lime- sulphur was first employed in this country as a sheep dip, but in 1886 its use against the San Jose scale began. It is now either purchased in the FIG. 106. A simple home-made arrangement for making lime-sulphur. commercial form or made at home. It may be prepared by heating together 80 pounds of high grade flowers of sulphur, 40 pounds of burned lime, and 50 gallons of water. A rather pure lime should be used, one that is low in mag- nesium, or otherwise there will be a waste of sulphur. The lime is first slaked in an iron kettle; and the sulphur is separately thoroughly mixed with & small amount of water and then added to the lime, after which the mass is boiled for 45 minutes. The mixture is best when boiled by passing steam through it. During the heating process, the lime com- bines with the sulphur forming calcium sulphides of variable composition. On standing, there results a yellow to orange colored solution, under which is an insoluble sludge. This 312 CHEMISTRY AND DAILY LIFE consists of undissolved sulphur, calcium sulphite, and calcium sulphate. During the boiling, the solution should con- stantly be kept up to the 50-gallon mark by adding water as it evaporates. When finished, the clear liquid should be strained off and tightly barreled, because it does not keep in contact with the air, which oxidizes it. For the purpose of exterminating scale insects the liquid is diluted to test 4.5 to 5.0 Baume. For summer work, the solution is diluted with water to test but 1 Baume. By this is meant that the strength of the solution must be so adjusted that a Baume hydrometer (Fig. 107) will sink to the marks mentioned before the liquid is used. It is believed that on the tree the calcium sul- phides decompose and leave free sulphur as the active agent. Kerosene in the form of kerosene emulsion is also often used as an insecticide. Kerosene is a compound of hydrogen and carbon and is prepared from crude petroleum (see Chapter IX). It kills insects, and when applied to pools of stagnant water, it suffocates the emerg- ing pupse of mosquitoes. When applied di- rectly to plants, it kills them. Kerosene cannot be diluted with water because it will not mix with the latter. For this reason kerosene must be mixed with some material that will carry the oil, as it were, and keep it from separating out from the mixture. For this purpose a soap is usually employed. Kerosene emulsion may be made by dissolving one pound of soap in 2.5 gallons of water, then adding 2.5 gallons of kerosene to the solution, and thoroughly mixing by pumping the entire mixture through a bucket sprayer. Diluted to from 20 to 30 FIG. 107. A Baume hy- drometer in use. POISONS FOR FARM AND ORCHARD PESTS 313 gallons, the mixture may be used against scale and other sucking insects. Tobacco decoctions can also be used as germicides; as such, their value depends upon the poisonous properties of the nicotine they contain. This alkaloid is soluble in water, and FIG. 108. Plant lice (aphis) on maple leaf. so hot water extracts of the stalk and waste tobacco are made, cooled, and then used as an insecticide. It is common prac- tice to burn tobacco in greenhouses to get rid of certain plant lice. Gaseous insecticides are often used against insects that are particularly difficult to attack. Of these substances hydro- cyanic acid gas is the most effective. Cyanides are deadly 314 CHEMISTRY AND DAILY LIFE poisons and should never be handled with the fingers (see Chapter IX). Hydrocyanic acid gas is made by treating potassium cyanide with sulphuric acid. Two ounces of concentrated sulphuric acid are carefully mixed with 4 ounces of water; this is placed in an earthenware vessel and then 1 ounce of potassium cyanide is added. This is the proper quantity for 100 cubic feet of space. The gas is a deadly poi- son and one breath of it may be fatal ; it should therefore by no means be inhaled. To retain the gas about the plant and make its action effective, it should be applied in tightly closed rooms or buildings, or under tents, as is done in the treatment of San Jose scale on the orange trees of California. The inclosure should afterwards be opened from the outside and thoroughly aired before it is entered by any one. Carbon bisulphide is a colorless, volatile liquid made from carbon and sulphur (see Chapter VII). The liquid is volatile and its vapors are fatal to insects. Being heavier than air, the vapors will settle through a mass of grain, and so they are very effective in killing grain weevils. A teaspoonful for each cubic foot of space, placed in a shallow dish upon the surface of the grain, will kill the insects present. The vapors settle down through the grain, and may later, after their work is done, be released by boring holes through the walls of the bin. Ants, moles, prairie dogs, and similar pests are exterminated by placing cotton saturated with carbon bisulphide in the heaps or runs, and covering tightly with earth. Carbon bisul- phide should never be brought near a flame, for it is even more dangerous than gasoline. Fungicides are materials for the destruction of certain parasitic plants called fungi. The chief fungicide in common use is Bordeaux mixture, which originated in France for the control of the downy mildew on grapes. It is made by sus- pending 4 pounds of copper sulphate (also called blue vitriol POISONS FOR FARM AND ORCHARD PESTS 315 or bluestone, see Chapter XI) in a gunny sack in about 15 gal- lons of water. The copper sulphate dissolves to a clear blue solution. Six pounds of lime are slaked and diluted to about FIG. 109. Students spraying an orchard. 15 gallons. The two solutions are then mixed and diluted to 50 gallons. This is the Bordeaux mixture. It is a sus- pension of rather insoluble salts of copper and lime of complex 316 CHEMISTRY AND DAILY LIFE composition. When applied to the foliage, the copper is slowly brought into solution by the action of the carbon dioxide of the air, and thus becomes active in destroying the pest. The lime should always be in excess in Bordeaux, and sufficient in quantity so that the mixture will turn red litmus paper blue. As long as copper sulphate is in excess, the reaction will be acid to litmus and the scorching of the foliage will follow. The potassium ferrocyanide test may also be used to detect an excess of copper sulphate in solu- tion ; for when Bordeaux mixture is filtered and a solution of potassium ferrocyanide is added to the filtrate, a brown precipitate of copper ferrocyanide forms if free copper sul- phate is present. Bordeaux mixture is a fungicide, and it alone has no effective poisonous action upon plant insects. To make it an insecticide, either Paris green or lead arsenate should be added. This is now frequently done in the propor- tion of 4 ounces to 50 gallons of the Bordeaux mixture. There are still other poisons that are sometimes used in the home and barn which must be briefly discussed. For- malin, or formaldehyde, is a powerful disinfectant; that is, it causes the complete destruction of disease germs. It is a product of the oxidation of wood alcohol, and is put on the market as a 38 to 40 per cent solution in water. For killing smut spores on grain, the seeds are immersed for ten minutes in a solution of one pint of the above 40 per cent solution to 20 gallons of water. It must be understood that formalde- hyde is never used as a spray for foliage. For other purposes of disinfecting, as, for example, for use in water closets or for disinfecting sick rooms, it has no superior. Mercuric chloride, corrosive sublimate, is another common disinfectant (see Chapter XI). It is put up in tablet forms with ammonium chloride to hasten its solution in water. One part in 1000 makes a sufficiently strong solution to kill POISONS FOR FARM AND ORCHARD PESTS 317 most germs. Corrosive sublimate is a stomachic poison, and forms insoluble compounds with proteins. Consequently an antidote for it is raw eggs or milk. Chloride of lime or bleaching powder (see Chapter VI) is a very common disinfecting and bleaching agent. It is a powder prepared by passing chlorine into slaked lime. It always has a strong odor of chlorine. It is unstable and slowly gives up its oxygen, leaving calcium chloride behind. The oxygen thus liberated is the active agent which destroys germs and other organic matter. Javelle water, which is also used for disinfecting purposes, is made by treating bleaching powder with washing soda and allowing the precipitate to settle. The clear solution is employed. Lime alone is one of the cheapest and most useful of the various disinfectants. It will destroy organic matter as well as bacteria, and therefore it is useful in disposing of the bodies of animals which have died of some disease; it also serves as a whitewash in barns and pens. Sixty parts of water to 100 parts of lime will make a good whitewash or " milk of lime " ; it must be fresh to be active, for air-slaked lime is the carbonate and has no germicidal power. When ap- plied with a spray pump, whitewash will enter the cracks in stables, etc., and be very effective. Whitewash, containing a little carbolic acid, is an especially good remedy in the poultry house for lice and vermin. The burning of sulphur is one of the oldest methods of dis- infecting ; but it is not very reliable. Sulphur dioxide de- stroys insects, vermin, and animal life in general, but it does not destroy the spores of bacteria. In order to act at all, sulphur dioxide requires the presence of moisture. Sulphur may be procured everywhere, and it has the further advan- tage that it is cheap. In fumigating with sulphur proceed as 318 CHEMISTRY AND DAILY LIFE follows : In a washtub put water to a depth of about four inches. In this, place several bricks or flat stones so that they will protrude above the water about two inches. Upon the bricks or stones, which should be in about the middle of the tub, place an old iron kettle. See that it rests securely on the bricks. Now place pulverized sulphur in the kettle, from 1 to 3 pounds, according to the size of the room to be fumigated. Open all the doors of closets, cupboards, etc., in the room. Also spread out the bedding. Now, having the washtub near the center of the room, insert into the heap of sulphur a piece of blotting paper about 3 by 8 inches, which has been soaked in alcohol. Wood alcohol or denatured alcohol or even strong rum or whisky will do. Mount this paper so that about 3 to 4 inches will be immersed in the powdered sulphur. Now with a match light the upper end of the paper. The alcohol will take fire readily and burn with a hot blue flame which will be communicated to the sulphur. As soon as the alcohol is burning well, leave the room and close the door tightly. After several hours, all the doors and windows may be opened so as to air the room well. It should not be inhabited till all odor of sulphur dioxide is gone. In cities the health officer will attend to the fumigation of rooms that have been occupied by patients having contagious diseases. Certain products that are prepared from coal tar are also good disinfectants. Among these may be mentioned car- bolic acid, cresol, creosote, and naphthaline. The latter is used in moth balls. Lysol, carboleum, germol, and zeno- leum are used in dips. These are all coal tar products, as are also the so-called fly removers which are sprayed upon animals to protect them from flies. Some of the advertised lice exterminators contain tobacco and sulphur as the active materials. These are not as ef- POISONS FOR FARM AND ORCHARD PESTS 319 fective as lime and carbolic acid. In mew of the fact that so many of these secretly compounded materials are fraudulent, ineffective, or very expensive, it is well to be sure of what the mixtures contain so as to be able to judge of their worth before purchasing them. It should be remembered that nature has provided good, efficient disinfecting agents, which, however, must be allowed to act freely. Sunlight is a powerful germ killer. If sun- light could be carried into every nook or corner of the house or the barn, there would seldom be any need of chemical dis- infectants. The germs of tuberculosis, diphtheria, and typhoid fever are killed by direct sunlight in six to eight hours. The diffused sunlight which reaches our rooms is not nearly so powerful as direct sunlight, and it usually requires several days for it to destroy germ life. Another common agent for killing bacteria is heat. It may be employed dry, as by bak- ing in an oven, or moist as by boiling with water or steam. Dry heat is not as effective in destroying germs as moist heat. All germs in clothing may be destroyed by boiling it for from five to ten minutes. Floors and walls may be dis- infected cheaply and efficiently with boiling water to which some lye has been added. QUESTIONS 1. Name two ways in which insects feed. How do the two classes of poisons used to exterminate insects act ? 2. What is the common poisonous ingredient of the stomachic poisons ? What is Paris green made of ? Give two ways for test- ing Paris green for purity. 3. Why should lead arsenate be purchased as a paste ? 4. How may lime-sulphur be prepared ? Why should a high- grade lime be used ? 5. For what purpose is kerosene emulsion used ? Why is soap used in its preparation ? 320 CHEMISTRY AND DAILY LIFE 6. What is the substance used in killing ants, gophers, etc. ? Why should flames be avoided when using this substance ? 7. What is Bordeaux mixture, and why should lime be used in excess in its preparation ? 8. What is formalin and when can it be used advantageously ? 9. Name the chemical elements in bleaching powder. State how it acts. 10. Why is whitewash a good material to be used in stables and hen-houses ? 11. What is the action of sunlight on germs ? CHAPTER XXII PRACTICAL EXPERIMENTS TO BE PERFORMED IN THE LABORATORY 1 LABORATORY MANIPULATION Cutting glass tubing. Place the tubing on the desk, and firmly draw the edge of a triangular file across it two or three times (Fig. 110) at the place where the glass is to be cut. Then take the tube in the hands and with the thumbs placed near together, just back of the scratch, gently pull the glass apart, at the same time pressing outward with the thumbs (Fig. 111). If the tube does not break, make a deeper scratch. If the tube is large, it may be necessary to file around the glass. The broken edges will be sharp and should be rounded by rotating them in the tip of a Bunsen burner. Bending glass tubing. Use an ordinary gas flame, or the luminous Bunsen burner spread out by a " wing top." Hold the tube in the flame, as shown in Fig. 112. Rotate the 1 Lists of apparatus and chemicals required for these experiments are given at the end of this chapter. The teacher will see to it that students use proper care in handling flames and dangerous substances. Y 321 FIG. 110. Cutting glass tubing. 322 CHEMISTRY AND DAILY LIFE tube and heat it until it begins to bend by its own weight; then remove it from the flame and care- fully bend the tube to the desired shape. If the tube is heated uniformly and not too highly, the bore of the tube will not contract or flatten at the bend. (See Fig. 113.) FIG. 111. Breaking glass tubing. In FIG. 112. Bending glass tubing. Fitting corks to glass tubing. Select a cork of suitable size for the test tube or flask to be used. Soften the cork by means of a cork press, or by roll- ing it between the desk and a block of wood. Select a cork borer slightly smaller than the tube to be inserted. Place the cork on the desk and bore through it, FlG> 113 '~ Good t ^ n g ad bends f glass PRACTICAL LABORATORY EXPERIMENTS 323 not by merely pressing, but by rotating the borer under slight pressure, as in using a gimlet. Keep the borer up- right and at right angles to the cork. If the hole is too small, a round file may be used to enlarge it. Pouring liquids from bottles. Remove the bottle from the shelf, lift the stopper with the fingers, as shown in Fig. 114, and pour into the Fig. 114. Pouring from a bottle. test tube. When the required amount of acid has been poured out, touch the neck of the bottle with the test tube to catch the drops on the edge and thus prevent them from streaking down the side of the bottle and FIG. 115. Pouring from a beaker. 324 CHEMISTRY AND DAILY LIFE on to the shelf or table. When pouring from one beaker into another, a glass rod, held as shown in Fig. 115, will pre- vent the liquid from running down the side of the beaker. Filtering. This ordinarily has for its purpose the separa- tion of a liquid from a solid. Place the funnel in the arm of a wooden stand. Fold the filter paper into halves and then at right angles into quarters (Fig. 116). The folded filter paper is opened so as to form a cone half of which is three thicknesses of paper and the remainder one thickness. Now fit the cone into a funnel of such a size that the edge of the paper does not quite reach the top. If the paper does not FIG. 116. Folding a filter. FIG. 117. A filter ready for use. fit, vary the folding slightly until it does. Now place the paper into the funnel and moisten it so that it adheres to the sides of the funnel. Press the paper gently with the fingers against the side of the funnel to remove any air bubbles. The filter is now ready for use (Fig. 117). PRACTICAL LABORATORY EXPERIMENTS 325 Heating liquids in test tubes. Fill the test tube only about one-third full. Hold it between the fingers and thumb, and constantly rotate it in the flame so as to apply heat uniformly. The heat should be applied to the upper end of the liquid and not to the bottom or to the glass above the liquid, as FIG. 118. Heating a liquid in a test tube. shown in Fig. 118. A test-tube holder or a band of paper may be used to hold the tube. In the following experiments the student should record his observations and conclusions neatly and carefully in a labora- tory notebook, which is to be submitted to the teacher for cor- rections and suggestions upon completion of each experi- ment. 326 CHEMISTRY AND DAILY LIFE EXPERIMENTS TO ACCOMPANY CHAPTER I 1. Chemical and physical changes. Examine a piece of rock candy. Notice its properties, such as its color, taste, hardness, and crystalline form. Describe its taste. (1) Grind a piece in a mortar. Has the taste changed? (2) Partly fill a beaker with water, and add a little of the powdered candy and stir with a glass rod. Taste the solu- tion. Does the sugar still exist? (3) Heat 5 grams of the powdered candy in a dry test tube, applying the heat gently at first, and noting carefully the changes. When no further change takes place, even at a much higher temperature, allow the tube to cool. Taste the residue. Will it dissolve in water? Notice its color. What does the new substance resemble? Has there been a chemical change? (4) Heat a platinum wire or a piece of porcelain, and then let it cool. Was there a change in this case ? Was it physical or chemical ? 2. Compounds and mixtures. (1) Note the visible prop- erties of a piece of sulphur. Test it with a magnet. Is it attracted ? Drop a piece of roll sulphur, as large as a pea, in a test tube and add 5 to 10 cc. of carbon bisulphide and shake. Use carbon bisulphide sparingly. Remember to have no flame in the room while working with it, for it is more dangerous than gasoline. Does the sulphur dis- solve? Pour off the liquid from the undissolved sulphur into a watch glass, and let it evaporate. Is there a residue on the watch glass ? Is it like sulphur ? (2) Examine some powdered iron. Test it with a magnet. See if it will dissolve in carbon bisulphide. (3) Stir together 3 grams of powdered sulphur and 5 grams of powdered iron. What is the color of the new powder ? Bring a magnet to some of the mixture. Does the iron still exist ? Put some of the powder in a test tube, add carbon bisulphide, shake, pour off the PRACTICAL LABORATORY EXPERIMENTS 327 liquid, and evaporate some of it on a watch glass. Does the sulphur still exist in the mixture ? 3. Chemical change. Heat the rest of the mixture of sulphur and iron from the above experiment in a test tube with a small flame. When the mass begins to glow like a red-hot coal, remove the tube from the flame. Lay the tube down to cool. There is probably a little melted sul- phur part way up the inside of the tube. Do not allow this to run down and mix with the substance at the bottom of the tube. When the tube is cool, break it open and ex- amine the substance in the bottom where the glow had been. See if any sulphur can be dissolved out of it with carbon bi- sulphide. Test the black material with a magnet. Can you detect either sulphur or iron in the new substance ? EXPERIMENTS TO ACCOMPANY CHAPTER II 4. Electrolysis of water. This experiment is to be done by the instructor. Set up the apparatus as shown in Fig. 1. The water used should be slightly acidulated with sulphuric acid, for pure distilled water will not conduct the electric cur- rent sufficiently well. Three or four dry cells will suffice. Note the difference in the volume of gas formed in the two arms. 5. Explosion of hydrogen and oxygen mixture. This ex- periment is to be done by the instructor, who will prepare oxyhydrogen gas by electrolysis of water, and then pass the mixture into the eudiometer tube over mercury. The ex- plosion is effected by means of the electric spark, the appara- tus for this purpose being arranged as shown in Fig. 2. 6. Salts in well water. Place some ordinary service or well water on a watch glass, and set it in a warm place to evaporate. The evaporation will proceed best by placing 328 CHEMISTRY AND DAILY LIFE the watch glass on a hot water bath. Note the residue left in the dish. 7. Preparation of distilled water. Set up an apparatus as shown in Fig. 3. Place from 200 to 300 cc. of water in the flask. Apply heat gently, and collect 100 cc. of the distilled water. Be sure that cold water is running through the con- denser. Observe the difference in the taste and odor of the water before and after distillation. Evaporate a few cubic centimeters of the distillate by gentle heat on a clean watch glass. Is there any residue left? How does this result compare with that of experiment 6? 8. Water in various substances. (1) Take some green leaves of any plant, also a carrot, or a potato to represent a root, and some wheat or corn as samples of seed. Weigh out 50 grams of each in weighed dishes of tinned iron or porce- lain. Then put the dishes in a warm place and cover each of them with a glass plate. Window glass will do. In a few minutes the plates will be dimmed over with water that has been driven out of the plant by heat. Save the materials for experiment 9. (2) Heat pieces of beef, old mortar, and epsom salts each separately in a large test tube and observe the water that condenses on the side of the tube (see Fig. 4). The pieces to be heated should be of about the size of a small hickory nut. 9. Amount of water in a plant. When it has been seen that the plant tissue in experiment 8 is losing water, remove the glass covers and put the dishes in the open air or on the water bath to dry. This will require a day or so. On re- weighing the dishes, it will be found that the contents have lost a good deal of water ; 80 to 90 per cent in the case of the green leaves. The roots lose nearly as much, while the seeds lose only from 10 to 15 per cent. Saw the material for ex- periment 10. PRACTICAL LABORATORY EXPERIMENTS 329 10. Ash in a plant. Take the residue from experiment 9 and heat it gently over a flame in a porcelain dish until everything possible has been burned off. There will remain a gray or white ash which may amount to from 2 to 5 per cent of the dry matter that was left after removing the water. In the plant ash only a few elements are found, but these are the same whatever the plant, or wherever it was grown. The elements in the ash are present as phosphates, sulphates, carbonates, chlorides, and silicates. The teacher will show the presence of calcium and phosphorus in the ash by simple qualitative tests. 11. Deposits of lime by boiling water. Boil for a few minutes about 200 cc. of hard well water or " White Rock" mineral water in a flask. After the water is cool, note any sediment of lime or turbidity of the water. Boiling has ex- pelled the carbon dioxide, which has held the carbonates of lime and magnesia in solution. When this carbon dioxide is driven off, the carbonates precipitate. The crust in tea- kettles and the scale in boilers are chiefly of this nature. These deposits interfere with good heating and steaming. 12. Hard and soft water. Place about 50 cc. of very hard water in a 200-cc. cylinder or flask. This water may be prepared, if necessary, by adding 0.1 or 0.2 gram of calcium chloride to 500 cc. of ordinary water. Add to this a measured quantity of soap solution, made as stated below. Mix well, and notice how much soap solution must be added before a permanent lather is formed. Also notice the precipitate of lime soap. Repeat the experiment, using rain water and tap water respectively. Which of these waters requires the most soap solution ? Hardness of water is due to the lime and magnesia salts in solution. To prepare the soap solution, scrape 10 grams of castile soap to fine shavings, and dissolve these in a liter of denatured alcohol ; then dilute with water 330 CHEMISTRY AND DAILY LIFE to 1350 cc. If the solution is not clear, filter it. Keep it in a tightly stoppered bottle. EXPERIMENTS TO ACCOMPANY CHAPTER III 13. Preparation of hydrogen. Place from 15 to 20 grams of granulated zinc in a flask arranged as shown in Fig. 119. Cover the zinc with 25 cc. of water. The thistle tube should dip below the water. Fill two or three cylinders with water for collection of the gas in the pneumatic trough. See FIG. 119. Hydrogen generator. that the stopper is tight, for hydrogen easily escapes. When all is ready, add, through the thistle tube, about 15 cc. of concentrated hydrochloric acid. Do not apply any heat. PRACTICAL LABORATORY EXPERIMENTS 331 Do not collect any gas until the generator has run for at least two minutes. Wrap a towel securely around the generator, and then collect a test tube full of the gas and light it. If it burns quietly without explosion, proceed to collect a cylinder or small bottle (100 cc.) of gas. (1) Introduce a burning splinter into the gas; does it support combustion? (2) When the generator has been running for some time five to six minutes disconnect at A, attach the blowpipe tip B, and light the gas. Explosions will occur if the air is not all displaced. Hold an inverted clean dry beaker over the flame ; note the production of water by the combustion. The hydrogen unites with the oxygen of the air to form water. Use your brass blowpipe as a jet when lighting the gas, and note the color of the flame. The hydrogen flame is colorless. 14. Action of sodium on water. In your evaporating dish place 20 cc. of water and then drop into it a piece of sodium half as large as a pea. After the action has entirely ceased, test the solution with red litmus paper. What com- pound is dissolved in the water ? 15. Preparation of oxygen. Set up the apparatus as shown in Fig. 120. Fill the pneumatic trough nearly full of water and place in it the free end of the delivery tube. Weigh out 5 grams each of potassium chlorate and manganese dioxide. Mix these on a paper and place the mixture in the test tube. Have the test tube perfectly dry inside and out. Fill several cylinders with water, cover these with glass plates, and invert them on the shelf of the trough, removing the plates. With a small flame, apply heat cautiously to the test tube. The burner should be held in the hand, so that the flame may be removed from time to time ; it should not be allowed to strike in one spot, otherwise the glass may crack. As soon as gas comes off freely, place the end of the delivery 332 CHEMISTRY AND DAILY LIFE tube so that the gas is collected in one of the cylinders. When the cylinder is filled, cover it with one of the glass plates while still under water. It can then be placed upright on the desk, while another cylinder is being filled. After thus filling two or three cylinders, remove the end of the delivery tube from the water, and then remove the flame. Do not remove the flame while the delivery tube is under water. If you FIG. 120. Making oxygen. do, a vacuum will be formed and the water will be forced into the hot tube and break it. (1) Light a splinter and place it for a moment in one of the cylinders ; remove it and ex- tinguish the flame, but while still glowing, thrust it into the cylinder again. (2) Put a small piece of sulphur, no larger than a grain of wheat, into a small iron spoon. Ignite the sulphur in a flame and thrust the burning material into a cylinder of oxygen. Record the results. (3) Heat the end of a piece of fine iron wire. Dip into it flowers of sulphur, light the sulphur, and plunge the wire into a cylinder of oxygen while the sulphur is still burning. What is the result ? (4) Heat a piece of charcoal in the air, and while it is still glow- ing drop it into a cylinder of oxygen. What is the result ? Now shake the contents of the cylinder with 5 cc. of water ; PRACTICAL LABORATORY EXPERIMENTS 333 then test the water with a piece of blue litmus paper. Also add clear limewater and shake the contents. Record the re- sults and explain your observations. 16. Ozone by electric sparks. Work a frictional electric machine and note the odor produced in the air. Ozone has been formed by the electric discharges. 17. Ozone. Smell of a bottle containing yellow phos- phorus in water. The odor of ozone can be detected. EXPERIMENTS TO ACCOMPANY CHAPTER IV 18. Carbon dioxide in the air. Pour 10 cc. of perfectly clear limewater (calcium hydroxide) into a test tube and blow air through it, using for the purpose a clean glass tube. Ob- serve the precipitate of calcium carbonate. Expose some clear limewater to the air in a beaker for 24 hours and ob- serve the result. There is a small amount of carbon dioxide in the air and this is the plant's source of carbon. 19. Preparation of nitrogen. Float a thin piece of cork on a surface of water in the pneumatic trough. Place on it a piece of phosphorus half as large as a pea. Note : Phos- phorus is very dangerous. Handle it with a forceps. Keep it under water till you actually require the piece for use. Use no more than directed. Ignite the phosphorus with the end of a hot wire, and quickly place over it a wide-mouthed bottle filled with air, being careful to keep the mouth of the bottle under water. Let the gas stand over water until the white fumes are absorbed. The white fumes are phos- phorus pentoxide. Note the extent of the rise of the water in the bottle. (1) Now test the gas with a lighted splinter. Does it support combustion? (2) Place some burning sulphur on a spoon and insert it in the nitrogen. How does it behave ? 334 CHEMISTRY AND DAILY LIFE' 20. Preparation of nitric acid. Special care should be exercised by the student in the preparation of all acids; but with nitric acid special care must be taken, because if any is spilled on the hands it causes painful burns and wounds that heal slowly. Arrange the apparatus as shown in Fig. 121. Carefully introduce 25 grams of sodium nitrate and then 25 FIG. 121. Making nitric acid. grams of concentrated sulphuric acid into the retort, and shake very gently until all is well mixed. Heat the mixture gently and collect the nitric acid in a test tube, cooled by water. Make the following tests : (1) Remove a drop of the acid by means of a glass tube and apply it to a piece of woolen cloth or silk ; observe the result. (2) Place a few drops of indigo solution in a test tube containing 5 cc. of water and then add 2 cc. of nitric acid. What happens ? (3) Place a small piece of copper in the test tube containing the remain- der of the acid. If no reaction takes place, add a little water. What becomes of the copper ? 21. Preparation of ammonia. Arrange apparatus as shown in Fig. 122. Into the flask place 10 grams of ammonium chloride, 10 grams of powdered lime, and 25 cc. of water. PRACTICAL LABORATORY EXPERIMENTS 335 When the apparatus is properly connected, ap- ply heat gently for 8 to 10 minutes. (1) Test the gas with red litmus paper. (2) Test the gas with a burning splinter. (3) Into an evaporating dish place 5 cc. of hy- drochloric acid and 10 cc. of water, and pour this into a cylinder of the gas. Avoid inhaling any of the ammonia gas. The white fumes are ammo- nium chloride. Why did you collect the ammonia gas in an inverted cylin- der? What is formed when ammonia gas is treated with Water? FIG. 122. Making ammonia. Shake a cylinder of am- monia gas with 25 cc. of water, and test the solution with red litmus paper. EXPERIMENTS TO ACCOMPANY CHAPTER V 22. Acids in fruits and silage. (1) Cut open a lemon and squeeze some of the juice into a beaker. Add a few drops of phenolphthalein solution and then carefully add a dilute solution of potash or baking soda until the red color just appears. Note the lowering of the acid taste as the potash is added. Lemons contain citric acid. This ex- 336 CHEMISTRY AND DAILY LIFE periment may be repeated with grapes in which the main acid is tartaric acid. From the grapes we get cream of tartar. (2) Press the juice from silage, and neutralize the acids in this in the same way. In silage the acids are mainly acetic and lactic acids. 23. Acids and bases. Preparation of a salt. Put 10 cc. of dilute hydrochloric acid and 10 cc. of water in a porce- lain dish. Measure out 10 cc. of sodium hydroxide solution into a beaker, and add 50 cc. of water. Add this dilute sodium hydroxide to the acid, a little at a time, until the solution is neutral to litmus. Do not drop the paper into the solution, but by means of a glass rod transfer a drop of the solution to the paper. In case too much sodium hydroxide has been used, add a drop or two of the acid. Bases or alkalies turn red litmus blue, while acids turn blue litmus red. When the solution is neutral, it has no perceptible action upon litmus paper. Place the dish upon a sand bath and apply heat until the solution is evaporated to dryness. Avoid excessive heating. This will prevent spattering when the solution becomes concentrated. Taste some of the material in the dish. It is common salt and entirely different in taste and chemical action from either hydrochloric acid or sodium hydroxide. It is formed by the union of an acid and a base. Write the equation. EXPERIMENTS TO ACCOMPANY CHAPTER VI 24. Preparation of hydrochloric acid. Into a generator, arranged as in Fig. 119, place 15 grams of common salt. Through the funnel tube add 25 cc. of concentrated sul- phuric acid. Warm the mixture gently if necessary. Let some of the gas bubble through water in a receiving flask. Test the liquid with litmus paper. Taste it carefully. Take PRACTICAL LABORATORY EXPERIMENTS 337 away the water, and collect some of the gas in a bottle by displacement of air. Is the gas heavier or lighter than air ? Does the gas burn? Hold a filter paper moistened with ammonia near the escaping gas. Explain, and write the equation. Invert a jar of the gas over water. What causes the water to rise ? To some of the solution add a few drops of silver nitrate. A precipitate is formed which is the very insoluble silver chloride. 25. Chlorides in water. To test for chlorides, add a few drops of nitric acid to 10 cc. of water and then add 5 cc. of a 1 per cent solution of silver nitrate. The formation of a white precipitate of silver chloride indicates chlorides. Try this on well water and on service water, also on a 0.1 per cent solu- tion of common salt. 26. Bleaching action of bleaching powder. Place 20 grams of bleaching powder in a beaker. Cover the powder with water, and then add 40 cc. of 10 per cent sulphuric acid. Hang a moist strip of red calico in the beaker by means of a wire. Now cover the beaker with a glass plate, warm gently, and let it stand. What is the result ? Turkey red calico will be almost unchanged, but other cheaper calicoes will be altered. 27. Vapors of iodine. Place 1 gram of potassium iodide and an equal amount of manganese dioxide in a test tube, and then add 2 cc. of concentrated sulphuric acid. Warm gently. Note the violet color of the vapors formed. These are the vapors of iodine. 28. Vapors of bromine. Mix 1 gram of pulverized po- tassium bromide and 1 gram of manganese dioxide. Place this mixture in a test tube, add 2 cc. of concentrated sul- phuric acid, and heat gently. Note the brown fumes of bromine which are evolved. 338 CHEMISTRY AND DAILY LIFE EXPERIMENTS TO ACCOMPANY CHAPTER VII 29. Properties of sulphur. Place 15 grams of roll sul- phur, coarsely pulverized, in a test tube attached to a test tube holder, and heat slowly until it is a thin, amber- colored liquid. As the temperature increases, notice that the liquid becomes darker, until it is quite dark and so thick and viscid that it cannot be poured. Continue to heat until the material becomes slightly lighter in color and again is a liquid. Then quickly pour it into water, and when cool examine the mass and describe its properties. On standing several days it reverts to ordinary sulphur. Is sulphur soluble in water? Try to dissolve some roll sulphur in carbon bisulphide. Pour off the clear liquid on a watch glass and allow it to stand and evaporate (no flame should be near). Examine the crystals under a lens. 30. Sulphur dioxide. Fill a small iron spoon half full of sulphur. Ignite the sulphur and then lower it into a small cylinder containing a piece of wet, colored cloth, a piece of red litmus paper, or a red carnation. As soon as the sulphur stops burning, remove the spoon and cover the cylinder with a glass plate. What is the effect on the cloth, on the litmus paper, and on the flower? Sulphur dioxide destroys organic coloring matter and is also a good disinfectant. It kills germ life. 31. Sulphuric acid. Make the following tests with some of the sulphuric acid used in the Babcock test for milk fat. (1) Put 2 or 3 cc. in a test tube and stick a splinter of wood into it and leave it there. Then remove the splinter. Wash off the acid, and examine the stick. It has turned black. What is this black material ? Also drop a piece of sugar, about a gram, into 2 cc. of the sulphuric acid. (2) Put 10 cc. of water and 1 cc. of sulphuric acid into a test tube ; then add PRACTICAL LABORATORY EXPERIMENTS 339 2 to 3 cc. of barium chloride solution. Observe the result. Barium sulphate has been formed. This is a very insoluble precipitate, and this test is a general one for sulphates. (3) Let the barium sulphate settle, pour off the clear supernatant liquid, and try to dissolve the white salt in hot hydrochloric acid. Does it dissolve? 32. Hydrogen sulphide. Arrange the generator as shown in Fig. 119. The delivery tube and cork should fit tightly, and the delivery tube should pass into a solution of lead nitrate. Set the apparatus under a hood. Place 10 grams of iron sulphide in the generator, add 30 cc. of dilute hydrochloric acid, and immediately connect with the delivery tube. Note the precipitate formed. Pass the gas into a solution of common salt, then into a solution of copper sul- phate, then into a solution of tartar emetic. Note the results. Note the odor of the gas. It is formed when organic matter decomposes, as in the rotting of eggs and manure. 33. Test for phosphates. In a 50-cc. beaker on a sand bath, dissolve half a gram of bone ash in 10 cc. of dilute nitric acid and 20 cc. of water. Filter, and to the filtrate add 5 cc. of ammonium molybdate solution and then stir the mixture. Observe the yellow precipitate, which is ammo- nium phosphomolybdate. This is a general test for phos- phates. It should always be made in a faintly acid solution. 34. Insoluble phosphates. Dissolve one-half gram of sodium phosphate in 10 cc. of distilled water and then add 10 cc. of water containing half a gram of calcium chloride in solution. Observe the result. The precipitate is calcium phosphate. Repeat the experiment, using alum instead of calcium chloride. In this case the precipitate is aluminum phosphate. 340 CHEMISTRY AND DAILY LIFE EXPERIMENTS TO ACCOMPANY CHAPTER VIII 35. Making boric acid from borax. Boil 100 cc. of water and stir in powdered borax till no more will dissolve. Then acidify with hydrochloric acid. On cooling, the flakes of boric acid separate out. Filter off and test the solid residue as follows: (1) Dissolve some of the boric acid in 5 cc. of water contained in a porcelain evaporating dish. Add 2 drops of sulphuric acid and 10 cc. of alcohol to the solution, and then carefully apply the flame to the dish. When the alcohol begins to evaporate, ignite it with the flame. It will burn with a green flame which shows the presence of boric acid. (2) Test some of the aqueous boric acid solu- tion with turmeric paper. Note the color produced. The result is better if the boric acid solution is acidified with hydrochloric acid and the turmeric paper is dried at 100 C. after having been moistened with this solution. 36. Insolubility of silicates. Weigh out carefully 5 grams of ordinary soil and place it in a flask and add 25 cc. of concentrated hydrochloric acid. Does it dissolve? Warm for an hour ; pour off the acid through a filter paper, and wash the rest of the soil on to the filter paper. Put the funnel containing the soil in a warm place and let dry. Then reweigh the residue after tapping it out of the paper, being careful that none is lost. Has the acid dissolved very much ? The soil is largely made up of silicates, which are almost insoluble materials. 37. Sodium silicate. Place 1 cc. of a sirupy solution of sodium water glass in an evaporating dish, dilute with 10 cc. of water, and add 2 to 3 cc. of concentrated hydrochloric acid. Note the gelatinous precipitate. Evaporate to dry- ness, and heat the dish gently over a low flame. Cool, add water, filter, and examine the solid residue. What is it ? PRACTICAL LABORATORY EXPERIMENTS 341 EXPERIMENTS TO ACCOMPANY CHAPTER IX 38. Preparation of carbon. Place 2 pieces of wood and a piece of bone, each as large as a bean, in an iron crucible and cover the materials with sand. Heat the crucible until smoking ceases. Remove from the flame, let cool, and ex- amine the charcoal and the bone black. What are the prin- cipal elements in wood? In bone? Try the same experi- ment with some ground wheat. Particles of carbon may be obtained from a luminous gas flame or a candle or lamp flame, by holding a piece of cold porcelain in the flame. Carbon is deposited in chimneys as soot when fuel is burned with a poor draft. When fuel is burned completely, the carbon disappears entirely as the gas, carbon dioxide. 39. Reducing power of carbon. Mix thoroughly from 2 to 3 grams of pulverized copper oxide and an equal bulk of pulverized charcoal. Place the mixture in a small test tube and apply heat. Observe the result. What is the bright red material produced? What causes the oxygen to be separated from the copper oxide ? 40. Absorbing power of carbon. To 10 cc. of water in a test tube add 5 drops of cochineal solution and 10 cc. of hydrogen sulphide water. Add 3 to 5 grams of bone black or animal charcoal. Stopper the tube with a cork and shake it. Let stand 5 minutes, and then pour upon a filter. Now smell of the liquid and look at the color of the fil- trate. It has lost its color, and is odorless. The charcoal absorbs both colors and odors. A half bushel of charcoal in a bag hung in the water of a cistern will gradually remove the odor of the cistern water. 41. Preparation of carbon dioxide. Arrange the appa- ratus as shown in Fig. 119. Put 20 grams of marble in the flask and sufficient water to cover the end of the thistle tube. 342 CHEMISTRY AND DAILY LIFE Fill 2 or 3 cylinders with water for collecting the gas. Then add slowly through the tube 30 cc. of concentrated hydro- chloric acid. Allow some of the fresh gas to escape, and then collect from 2 to 3 cylinders of the gas. Remove the cylin- ders, and set them upright on the table. (1) While the ap- paratus is running pass some of the gas into a test tube of limewater, and note the result. A precipitate is first formed and gradually then redissolved. Boil the clear solution thus formed and note the reappearance of the precipitate. (2) Test some of the escaping gas with a burning splinter. (3) Pour a cylinder of the gas over a candle flame. What happens? (4) Invert another cylinder over a 25 per cent solution of sodium hydroxide. Why does the gas dis- appear ? Agricultural students may also perform similar experi- ments, using other calcium-containing substances, such as caustic lime, limestone, floats, and gypsum, to see if these contain carbonic acid, as does marble. 42. Making hard soap. In a beaker dissolve 15 grams of caustic soda in 120 cc. of water. Pour half of the water into a porcelain dish of about 500 cc. capacity, add 50 cc. of water and 50 grams of tallow. Boil the solution for 45 minutes, carefully replacing from time to time the water lost by evaporation. Then add the remainder of the caustic soda and boil for at least an hour more. The volume of the liquid may finally be allowed to decrease about one-third. Add 20 grams of common salt, boil for a few minutes, and allow to cool. The soap rises to the top, and the glycerine, excess of lye, and salt remain in solution. Try to make soap similarly, using gas engine cylinder oil instead of tallow. 43. Alcoholic fermentation. Weigh 10 grams of flour into a 500-cc. bottle, add 50 cc. of water and a small piece PRACTICAL LABORATORY EXPERIMENTS 343 FIG. 123. Alcoholic fermentation. of a yeast cake. By means of a delivery tube connect the flask with a test tube containing enough clear limewater to cover the end of the tube (see Fig. 123). Cover the lime- water with one- quarter of an inch of kerosene. Why? Place the bottle on a warm sand bath (not over 85 F.) for half an hour. Why? Observe the bubbles of gas given off and the precipitate that is formed in the lime- water. Carbon dioxide and alcohol are formed. This is what occurs in bread making. 44. Removal of grease spots from clothing. Spots due to fats such as butter, tallow, olive oil, etc., can be removed with soap or with gasoline or benzine ; also by placing the fabric between blotting paper and applying a warm iron. Try the solubility of fats in gasoline or benzine, also the removal of a grease spot by the hot iron method. Grease may also be removed by ammonia water or even a solution of borax. In the case of spots due to mineral oils, the spot should first be treated, or softened with some oil and then treated with gasoline. Why? Gasoline is inflammable. Do not have flames around when you are using it. 45. Washing soda. This is sodium carbonate and is sometimes used as a filler for soap. It is the main constitu- ent of washing powders. To test such a powder for sodium carbonate, put 5 grams in a test tube and add a few drops 344 CHEMISTRY AND DAILY LIFE of hydrochloric acid. If there is a brisk evolution of gas, it indicates the presence of a carbonate. 46. Odors and appearance of some important organic compounds. The instructor should have the student ex- amine with the eye and nose the following materials : car- bolic acid, ether, banana oil, carbon tetrachloride, chloro- form, etc. 47. Test for nitrogenous organic compounds. Mix a gram of dry clover hay, peas, or meat with enough soda- lime to half fill a test tube. Connect the test tube with a delivery tube as in Fig. 120, but let the end of the tube lead into another test tube containing water. Apply heat for from 5 to 10 minutes. Test the water with red litmus paper. It should turn blue. Soda-lime decomposes protein material and liberates nitrogen as ammonia. The carbon and oxygen of the protein unite and form carbon dioxide and water; the carbon dioxide forms carbonates with the soda- lime. 48. Test for proteins. Make about a 10 per cent caustic potash solution by dissolving 5 grams of caustic potash in 45 cc. of water. In another glass dissolve a piece of copper sulphate (bluestone) the size of a pea in 50 cc. of water. Pour about one-fifth of the white of an egg into 50 cc. of water. Place 10 cc. of this egg solution in a test tube, add 5 cc. of the potash solution, put the thumb over the tube, and shake until all is well mixed. Now add a few drops of the copper sulphate solution and shake again. At first the only color will be a greenish blue, but on standing this will change to a deep violet color. This is the test for proteins. Test wheat, corn, bran, and meat by this means. These materials should be crushed or ground, and warmed for a short time in a little of the caustic potash solution and then filtered before adding the copper sulphate solution. PRACTICAL LABORATORY EXPERIMENTS 345 49. Preparation of a protein. Put about 25 grams of flour in a porcelain dish, add 12 to 15 cc. of water, and mix into a stiff dough. Do not have too much water present. Let stand one-half hour, remove the ball of dough, and, hold- ing it in the fingers, add water little by little and so wash out the starch by further manipulation with the fingers. Continue the washing until the liquid runs clear, and then place the ball of gluten in water for one-half hour. Remove it from the water, freeing it from water as much as possible, and then weigh it. Spread it out in a thin cake, and then let it dry in the oven or in a warm place, and finally reweigh it. Calculate the per cent of gluten in the flour. The gluten consists of the two principal proteins of the wheat. 50. Fats and oils. Crush 5 grams of flax or hemp seed on a piece of white paper. A grease spot will appear on the latter. Should the test fail to show oil, place the crushed seeds and the paper in a warm place on a tin plate. The higher temperature will melt the oil out, which will then be shown by the grease spot on the paper. 51. Fats and oils. Crush four or five kernels of corn. Place the powder in a flask, and pour over it 25 cc. of gasoline. Shake it at intervals for 15 minutes. Keep the flask loosely stoppered and away from flames. Then pour the gasoline off, and through a filter paper, catching the filtrate in a beaker. Allow the gasoline to evaporate slowly. A resi- due is left in the beaker. Test it between the fingers for its greasy feeling. Gasoline and benzine dissolve fats. Grease spots on cloth are removed by washing with gaso- line or benzine. 52. Preparation of starch. Make a pulp of two clean potatoes. Tie the pulp in a cloth and squeeze the juice into a beaker filled with water, occasionally dipping the bag into water. Allow the beaker to stand for 20 minutes, or 346 CHEMISTRY AND DAILY LIFE until* the starch has settled. Now pour off the water, and if the starch is not clean, wash it by adding more water and again allowing it to settle. Finally pour off the water. Leave the beaker in the desk until the starch is dry, and then save the starch for the following tests. 53. Test for starch. Taste some of the starch. It is practically tasteless. Place one-half gram in a test tube about half full of water. Shake the test tube, then boil, and filter. To the filtrate add a few drops of iodine solution. (This is made by dissolving 5 grams of potassium iodide in 50 cc. of water and then adding 1 gram of iodine to the solution.) A deep blue color appears throughout the solution. Treat some of the starch with cold water, but do not filter ; add some of the iodine solution. On standing, the starch, settling out at the bottom, will be blue. In the first* case the starch was in solution, in the second it was not. Prep- arations of starch for the iodine test may also be made from corn, wheat, oats, or beans. 54. Test for sugar (glucose). Grind 25 grams of dried raisins in 50 cc. of water. Filter the solution into a beaker. Into a test tube place 10 cc. of Fehling's solution. This is made by dissolving 6.2 grams of copper sulphate, 3.5 grams of Rochelle salts (sodium potassium tartrate) , and 2 grams of potassium hydroxide in 100 cc. of water. Place about 10 cc. of the Fehling's solution in a test tube and add from 5 to 10 cc. of the raisin solution. Heat over a flame and let the liquid boil a few seconds. A reddish brown precipitate forms, which on standing deposits on the bottom of the test tube. This precipitate is the red oxide of copper, Cu 2 O, and shows the presence of glucose in the solution. Test the water extract from corn or wheat for glucose ; also the extract of an apple, a sample of honey, and one of "Karo sirup." PRACTICAL LABORATORY EXPERIMENTS 347 EXPERIMENTS TO ACCOMPANY CHAPTER X 55. Preparation of potassium hydroxide. Dissolve 5 grams of potassium carbonate or wood ashes in 15 cc. of water in an evaporating dish. Add a mixture of 3 grams of barium hydroxide (also called barium hydrate) and 10 cc. of water. Heat on a sand bath for 5 minutes. Filter off the solution and observe the precipitate. Evaporate some of the filtrate to dry ness by heating on the sand bath. This is potassium hydroxide. Keep it for experiment 56. Wood ashes contain potassium carbonate and are often the source of the " lye " used in soap making on the farm. 56. Test for potassium. Dip a clean platinum wire into a solution of the residue left from the last experiment. Hold it in a non-luminous Bunsen flame and observe the flame through a blue glass. A violet color shows the presence of potassium. If no platinum wire is at hand, use a clean strip of asbestos paper. Hold a crystal of potassium chlorate in the flame with the forceps. 57. Test for sodium. Dip a clean platinum wire into a solution of common salt; hold it in the Bunsen flame and observe the bright yellow color imparted to the flame. This is a test for sodium. Try the same test by holding a piece of rock salt in the flame with a forceps. 58. Precipitation of barium sulphate. To a solution of barium chloride add a drop of sulphuric acid. Note the heavy white precipitate formed, which is the insoluble barium sulphate. To a solution of copper sulphate add a few drops of barium chloride. Here again the insoluble barium sul- phate is formed. 59. Tests for the alkaline earth metals. Using a clean forceps, hold a piece of a salt of strontium in the flame and note the color produced. Perform the same experiment, 348 CHEMISTRY AND DAILY LIFE holding a piece of a calcium salt in the flame ; also repeat the experiment with a piece of a barium salt. 60. Testing the quality of lime. Place about 40 grams of burned lime in an evaporating dish, and moisten it well with water, warming to about 35 C. to hasten the action, if necessary. Note the reaction and observe how the mass warms up. Good, fresh lime readily undergoes the slaking process. Place some of the slaked lime in a bottle; add 100 cc. of distilled water ; shake vigorously, and allow to stand for four or five hours. Then filter some of the solution and test it by blowing air through it, using a glass tube. Place about one-half gram of the slaked lime in a test tube, add 10 cc. of water and then a few drops of hydrochloric acid. When the action ceases, add more acid, a little at a time, and warm. What has not dissolved consists of silica and clay. Lime of high purity contains less than 10 per cent of acid- insoluble material. 61. Plaster of Paris. In a test tube carefully heat 10 grams of powdered gypsum to 115 C. Regulate the temperature by means of a thermometer so that it reaches no higher than 115 C. W T hy? Allow it to cool. Mix with a small quantity of water and note that in a few min- utes the mass becomes hard. Explain. 62. Burning magnesium. Hold a piece of magnesium ribbon about an inch long in a forceps and apply a lighted match. The ribbon burns with a brilliant white flame. Examine the white product formed. What is it? Flash- light powder consists of about 5 parts of powdered magnesium to 9 parts of potassium chlorate. It should be used with care. PRACTICAL LABORATORY EXPERIMENTS 349 EXPERIMENTS TO ACCOMPANY CHAPTER XI 63. Testing for alum. Add a few drops of an alum solu- tion to a test tube containing 5 cc. of water and then add a few drops of logwood extract and 2 cc. of ammonium car- bonate solution. Observe the result. Mix about 2 grams of flour in a dish with water containing a few drops of alum solution. Add a few drops of logwood extract and the same amount of ammonium carbonate solution. Mix well and observe the result. Repeat the test, using a baking powder, testing for the presence or absence of alum. In the presence of alum a blue color is always obtained with tincture of log- wood and ammonium carbonate solution. 64. Compounds of alum and protein. To 5 cc. of a solu- tion of egg white add a few drops of an alum solution and observe the result. A heavy precipitate will form. This is an insoluble compound of protein and alum. To a solu- tion of aluminum chloride add ammonia water. The same result is noticed, but in this case aluminum hydroxide is formed. 65. Copper in a silver coin. Under a hood, dissolve about one-fourth of a dime in dilute nitric acid. It is best to hammer out the coin and then cut it. Account for the bluish color of the solution. To the solution add a solution of hydrochloric acid until no further precipitate is formed. The white precipitate is silver chloride. Filter and evapo- rate the blue solution of copper nitrate in an evaporating dish, finally igniting to redness over a flame. The black residue is the oxide of copper. Expose the silver chloride to the light. What happens? 66. Copper and its reactions. (1) Dissolve 5 grams of copper sulphate in 100 cc. of water. Place a bright iron nail in the solution and note the result. (2) To 5 grams of 350 CHEMISTRY AND DAILY LIFE Rochelle salts add 10 cc. of water, and 5 cc. of concentrated caustic soda solution, after all the salt has dissolved. Now add 5 cc. of the copper sulphate solution and then treat the resulting solution with 1 cc. of glucose (" Karo sirup ") and boil. A reddish brown precipitate of cuprous oxide, Cu 2 0, is formed. This is the test for sugar. (3) To 10 cc. of the copper sul- phate solution add a few drops of potassium ferrocyanide. A reddish brown precipitate is formed. This is a delicate test for copper, and can be used for testing Bordeaux mix- ture for excess of copper sulphate in solution. In Bordeaux preparation, if this test shows the brown precipitate, more lime must be added ; see experiment 151. 67. Copper in brass. Under the hood, treat about 0.2 gram of brass with 5 cc. of nitric acid in a test tube and note the blue color of the solution. This color is due to the for- mation of copper nitrate. 68. An antidote for mercury poisoning. To 20 cc. of water add 5 or 6 drops of the white of an egg. Stir thor- oughly and then add to this solution a few cubic centimeters of a dilute solution of mercuric chloride (corrosive sublimate), 1 in 2000. The precipitate is a compound of the protein and mercury. In case of mercury poisoning the whites of eggs and milk are a good antidote. 69. Preparation of mercury. To 5 cc. of a solution of stannous chloride add a few drops of a solution of mercuric chloride. The gray precipitate that separates is metallic mercury. Note the physical properties of some mercury contained in a bottle on the reagent shelf. 70. Lead and its salts. Dissolve 5 grams of metallic lead in warm, dilute nitric acid. Evaporate off nearly all of the excess of acid and allow the liquid to stand and cool. Pour off the remaining liquid from the crystals and dissolve these in 50 cc. of water. Test separate portions of 5 cc. each of this PRACTICAL LABORATORY EXPERIMENTS 351 solution with (a) hydrogen sulphide; (6) dilute sulphuric acid; and (c) dilute hydrochloric acid. Precipitates of dif- ferent colors are formed. Lead sulphide is black; lead sulphate and chloride are white. 71. Action of water on lead. Put a gram of clean bright lead shavings into a test tube containing 10 cc. of distilled water. After 24 hours decant the clear liquid into a second tube. Slightly acidify with hydrochloric acid and then add 5 to 10 cc. of fresh hydrogen sulphide water. A black or brown coloration indicates lead in solution. Soft water, especially peaty water, attacks and dissolves lead more than does hard water. In contact with hard water the lead be- comes coated with the insoluble lead sulphate and carbonate of lime which protect the metallic surface from further attack. 72. Preparation of lead chromate. To 5 cc. of a solution of potassium dichromate add 3 cc. of lead acetate (sugar of lead) solution. The beautiful yellow precipitate formed is chrome yellow, or lead chromate, which is used as a pigment in paints. 73. Making formalin. Treat 10 to 20 cc. of wood alcohol with one-half gram of potassium permanganate and 1 to 2 drops of strong sulphuric acid. Heat the beaker gently. Note the odor produced. The penetrating smell is due to the production of formaldehyde, formed from the wood alcohol by oxidation of the latter by the permanganate. 74. Cobalt bead. In a loop of platinum wire, made by bending the wire around the sharp end of a pencil, place some borax and then fuse it in the flame. Note the color of the bead. Now dip the bead into a dilute solution of cobalt nitrate. Then put the loop with the bead into the flame again and allow the bead to melt. Note the beautiful blue color of the bead on cooling. 352 CHEMISTRY AND DAILY LIFE 75. Preparation of iron sulphate. Place 25 grams of iron nails in a 500-cc. flask, add about 25 cc. of dilute sulphuric acid (1 part acid to 5 parts water), warm gently until no more gas comes off. Pour the clear solution into a clean beaker. Add 2 to 3 cc. of dilute sulphuric acid to the solu- tion, boil the liquid down to one-half of its volume, and allow it to cool. Crystals of iron sulphate will separate out. The liquid should be poured off and the crystals allowed to drain. This is the material used in the destruction of weeds and is commonly prepared at steel plate mills. Other names for this salt are copperas, green vitriol, and ferrous sulphate. 76. Change of ferrous to ferric iron. Dissolve one-half gram of ferrous sulphate in 20 cc. of water. Divide the liquid into two equal portions. To the first portion add a few drops of ammonia water until a precipitate is obtained. To the other portion add five drops of concentrated nitric acid. Heat this latter portion to boiling; cool, and then add am- monia water until a precipitate forms. Compare the two precipitates. The nitric acid oxidizes the iron, which then is in a substance containing more oxygen. The copperas is fer- rous sulphate ; by oxidizing it with nitric acid it was changed to ferric sulphate. 77. Compounds of tannin and iron (ink). (1) Dissolve 1 gram of tannic acid in 25 cc. of hot water. Dip a piece of cotton cloth (about 3 by 10 cm.) into the solution. Dry the cloth, and then dip it into a solution of ferrous sulphate (1 gram in 25 cc. of water). After the cloth has dried, see if the color can be removed by washing. (2) Add 5 cc. of ferrous sulphate solution to 5 cc. of tannic acid solu- tion. Observe the result. Does it look like ink ? 78. Making a blue print. Place a leaf or a drawing made on thin paper on the glass of a photographer's printing frame and then cover it with a piece of blue printing paper (the PRACTICAL LABORATORY EXPERIMENTS 353 sensitive side next to the leaf). This paper can be secured at most drug stores. The operation of placing the object and paper together should be done in a darkened room. Now expose the printing frame to direct sunlight for about 3 to 5 minutes. Several trials must be made to determine the time of exposure necessary for a successful print. Remove the paper from the frame. Wash with water and spread the print on a cloth to dry. EXPERIMENTS TO ACCOMPANY CHAPTER XII 79. Drying of linseed oil. With a brush spread a layer of linseed oil (very thin) on a well-varnished ruler and allow it to stand several days. Now examine it. Has it set to a hard film? Crush 10 grams of flaxseed in a mortar. Ex- tract the material with gasoline in a flask. Pour off the gasoline, and out of doors, or under a hood, allow the gaso- line to evaporate off. Paint the ruler with the oil that is left. Try a film of corn oil on a ruler. How does it act? 80. Solubility of oils. Try to dissolve some linseed oil in gasoline; also in carbon bisulphide. Try the same experi- ment, using corn oil, also butter fat, instead of linseed oil. EXPERIMENTS TO ACCOMPANY CHAPTER XIII 81. Odor of burning wool and cotton. Burn small pieces of wool, silk, and horse's hoof and note the odor in each case. Try the same experiment with cotton. The former materials are nitrogenous and give rise to ammonia and other odorifer- ous nitrogenous compounds, while cotton does not. Try the solubility of the various materials mentioned above in am- monia water. Wool, silk, and horn will dissolve, while cot- ton will not. 2 A 354 CHEMISTRY AND DAILY LIFE 82. Bluing. Indigo was very generally used for bluing in the laundry. Another material is now commonly employed ; namely, Prussian blue. This is ferric ferrocyanide, a complex cyanide of iron. Because of the ease with which it decom- poses with alkalies, there is danger that iron rust may be deposited on the goods if this form of bluing is used. To 10 cc. of a 10 per cent solution of ferric chloride, add 5 cc. of a 10 per cent solution of potassium ferrocyanide, and about 5 drops of hydrochloric acid. The blue precipitate is Prus- sian blue. Now add an excess of caustic alkali and heat to boiling. Note the reddish brown precipitate of ferric hy- droxide. This is what causes the iron rust stains. 83. Dyeing cloth. Dip strips of white woolen, silk, and cotton cloth into a solution of fuchsine. After immersion in the solution remove the strips and allow them to drain. When they are nearly .dry they may be washed in water to test the fastness of the color on the various fabrics. 84. Gelatine in sole leather. A piece of sole leather, half as large as the hand, should be boiled in water for an hour and then allowed to cool. Does the solution gelatinize or become viscid and sticky? Sole leather has been treated in such a way that it contains no gelatine. Try the same experiment, using a piece of fresh cowhide. Note the result. 85. Solubility of rubber. Place about 0.5 gram of rubber in a test tube, add 5 cc. of carbon bisulphide, heat gently in a hot water bath, and shake the contents of the tube. What happens? Notice how the rubber swells up and slowly dissolves. No flame should be near. 86. Solubility of asphalt. Dissolve some asphalt paint in turpentine and apply it to a piece of unpainted wood. Set the wood aside to dry and note the stain left behind. 87. Removal of iron rust from cloth. Stretch the stained cloth over -a dish containing hot water. Then as the steam PRACTICAL LABORATORY EXPERIMENTS 355 rises and the cloth becomes moist, drop a few drops of dilute (10 per cent) hydrochloric acid or, better, 10 per cent oxalic acid upon the rust spot, preferably with a medicine dropper. When oxalic is used the fabric may be immersed directly in the acid. After a moment, lower the fabric into the water. If the spot is not removed, repeat the operation. Then rinse in clear water and finally in dilute ammonia water to neutral- ize any remaining acid. Then rinse again. Iron rust stains may often be completely removed from fabrics by means of lemon juice. 88. Removal of paint stains. These can be removed by the use of turpentine. Spot some cloth with paint and, after it is thoroughly dried, remove the stain with turpentine. 89. Removal of ink stains. A great many ink stains can be removed with sour milk or with some dilute acid such as oxalic or citric acid. If the ink is an iron tannate, the methods of removing iron rust will be applicable. Stain some cloth with ink, and then remove the spot with the reagents men- tioned. These should be used in dilute solution (5 to 10 per cent) ; and after application and removal of the stain the fabric should be well washed with water. Long contact with the acid may weaken the fabric. 90. Removal of fruit stains. Fruit stains may be re- moved from many fabrics by washing in water containing a little borax or ammonia. Always try this treatment first, as it will not injure the colors. If the above does not remove the stain, try a very dilute solution of chloride of lime con- taining a few drops of dilute acetic acid (vinegar). When the stains have been removed, wash the fabric thoroughly with clean water. In the case of wool or silk, treatment with soap (care) and water is the only way. If not effective, the fabric must be re-dyed. Try these tests, and also the action of ammonia on wool and silk. 356 CHEMISTRY AND DAILY LIFE 91. Coffee and tea stains. These are removed in the same manner as the fruit stains. They are due to the color- ing matter in the liquids. Grass stains owe their color to chlorophyll. This is soluble in ammonia water or in alcohol, and can therefore be removed from fabrics by these re- agents. Apply these reagents to such stains. Also try leaving the stained material in the sunlight. What is the result? 92. Blood stains. These are due to the red pigment of the blood called haemoglobin. This coloring matter is easily decomposed by acids. It is also soluble in warm water, which should be tried for the removal of the stain, before using the acid treatment. EXPERIMENTS TO ACCOMPANY CHAPTER XIV 93. Nitrogen in soil. Mix 5 grams of soil and an equal bulk of soda-lime in a mortar and transfer the mixture to a strong test tube. Connect the tube with a delivery tube as in Fig. 120, but have the end of the tube lead into another test tube containing distilled water. Heat the first test tube cautiously with a burner for five to ten minutes. Test the water with litmus paper and note the reaction. The soda- lime, aided by heat, decomposes the organic matter of the soil and forms water, carbon dioxide, and ammonia. The latter contains the nitrogen and has passed over into the water. 94. Phosphorus in soils. Place 10 grams of soil in a flask, add 25 cc. strong hydrochloric acid, and warm on a water bath for half an hour. Pour off the acid, cool, and add am- monia gradually, shaking after each addition, and continue this until the solution turns red litmus blue. Then add nitric acid drop by drop until the solution is slightly acid, PRACTICAL LABORATORY EXPERIMENTS 357 or until the blue litmus becomes red. Then add 10 cc. of ammonium molybdate solution. A yellow color or precipi- tate after a few minutes shows the presence of phosphorus. It is best to keep the solution at 65 C. to aid the precipita- tion of the ammonium phosphomolybdate. 95. Reaction of soils. For this experiment use (1) peat, (2) a mildly alkaline soil, and (3) a clay soil. Bring in con- tact with each soil, moistened with distilled water, pieces of sensitive red and blue litmus paper. Note any changes in color of the litmus paper and record what results you find. In a similar way test the soils from the fields of your own farm. 96. Correcting the acidity. When you find an acid soil, - that is one which changes blue litmus to red, take a wash basin full of it and add 15 to 25 grams of burned lime or ground limestone. Stir the mixture thoroughly and test it again. Try the same experiment, using wood ashes instead of lime. If the soil is still acid, add enough lime un- til it shows a blue color with litmus. An alkaline or " sweet soil " is the normal one for most agricultural plants. 97. Weathering of limestone. Place about 25 grams each of fresh limestone, rotten limestone, and residual lime- stone soil in separate beakers. Add 100 cc. of 5 per cent hydrochloric acid to each, and allow them to stand for one hour. Filter off the insoluble material on filter papers al- ready dried and of known weight. Wash the residues with water. Dry them in the oven at 100 C. and weigh. Cal- culate the per cent of insoluble matter in each case and ex- plain the results. 98. Puddling of clay. In each of two basins place a pound of dry clay soil. To one of these add water until the soil is slightly moist, but do not work or stir it. To the other add enough water to make the clay sticky; mix thoroughly 358 CHEMISTRY AND DAILY LIFE with a stick and place the mixture in the sun to dry. Note the results. 99. Effect of soluble salts on soils. To 20 pounds of soil in a box, add 25 grams of sodium nitrate and mix thor- oughly. To another box of the same soil add 2 grams of the same salt and mix. Water both, and set them in a proper place for growth. Plant with the same number of seeds (25) and after the plants are well started thin to the same number (12). Note the results. Too much of soluble salts may be harmful. An acre of soil 6 inches deep weighs about 2,000,000 pounds; estimate the application per acre on the above amounts of nitrate used for 20 pounds of soil. Does this suggest why soluble fertiliz- ers must not be applied in large amounts ? 100. Effect of deep plant- ing on the germination of seeds. Place 2 to 3 inches of moist soil on the bottom of a glass jar. Plant 4 to 5 peas near the wall so that they may be observed. Now cover with 2 inches more of soil and plant some more seed. Re- peat this .(Fig. 124), planting the seeds only an inch deep. Keep the soil moist and warm. Record what happens. 1 01. Capillary action of water on soil. Firmly tie a piece of linen cloth over the ends of several glass tubes, one-half to three-fourths inches in diameter. Lamp chimneys are suitable. Fill one tube FIG. 124. Effect of deep planting on germination of seed. PRACTICAL LABORATORY EXPERIMENTS 359 with sandy soil, another with clay soil, and a third with a loam. Compact the soils after each addition by gently jarring. Then immerse the lower ends of the tubes in a basin of water. Support them upright and allow them to stand. For one week, measure each day the height to which the water rises. What does this show? 102. Sedimentation of clay. In each of three separate cylinders, or beakers, place 200 cc. of turbid water made by shaking up some clay. To beaker No. 1 add one-half gram of burned lime and stir; to No. 2 add one gram of burned lime and stir; the third beaker is used for purposes of com- parison, and no lime is added to it. After 24 hours observe the beakers, and note the influence of the lime in throwing down the clay and clarifying the liquid. Liming clay soils makes them work better. 103. Effect of color on soil temperature. Fill two boxes, each one foot square and five inches deep, with soil and gently compact this. Cover one with lampblack and the other with gypsum or chalk. Insert thermometers in the soil to a depth of 1.5 inches and expose the boxes equally to the rays of the sun, recording the reading of the thermom- eters every ten minutes for two hours. Arrange the results in a table. What is the conclusion? 104. Effect of drainage on soil temperature. Fill two boxes as in the above experiment, but have one box lined with tin or galvanized iron or oiled cloth. Add water to the soil in the box not lined until it begins to drain. Add the same amount to the other. Let stand until the water in the unlined box has largely drained away. Take temperature readings as in experiment 103. 360 CHEMISTRY AND DAILY LIFE EXPERIMENTS TO ACCOMPANY CHAPTER XV 105. Action of fertilizers. Into each of two boxes, about one foot square and five inches deep, place 20 pounds of or- dinary dry soil. Before putting the soil into one of the boxes, add to it the following materials, mixing them thor- oughly with the soil : 5 grams of calcium acid phosphate, 3 grams of potassium nitrate, and 3 grams of potassium sul- phate. Water the soils, and when they have been properly worked, plant in each box 25 rape seeds and place the boxes in a greenhouse or window. When the plants are up, thin them down to 16 and let them grow. Note the difference in growth from week to week. Try the same experiment, using sand as the soil, and note the results. What elements needed by the plants may be lacking in the sand, and are not supplied by the fertilizer you have used? Action on acid soil: try the same experiment, using an acid soil and clover seed instead of rape seed. To one of the boxes add 10 grams of finely ground limestone. Why add the lime- stone ? 1 06. Phosphorus in bones. (1) To about one gram of bone ash in 15 cc. of water add 3 to 5 cc. of nitric acid, shake the mixture and filter. To the warm filtrate add 5 to 10 cc. of ammonium molybdate, and warm. What is the yellow precipitate? What does it indicate? (2) In a test tube heat one gram of bone ash with 20 cc. of distilled water and filter. To the warm filtrate add 5 cc. of ammonium molyb- date solution. Note the result, and compare it with that obtained in the previous experiment (1) in which nitric acid was first added to the ash. Try the same experiment with acid phosphate ; also with floats. 107. Solubility of potash fertilizers. Place 1 gram of kainit in a test tube, add 100 cc. of water, and shake the mix- PRACTICAL LABORATORY EXPERIMENTS 361 ture. Does the salt dissolve? Try the same experiment with saltpeter (potassium nitrate) ; also with muriate of potash. These salts are all readily soluble potash ferti- lizers. 1 08. Making acid phosphate. In a cigar box mix 100 grams of bone meal and 100 grams of commercial sulphuric acid. Add the acid slowly, stirring constantly with an iron rod. After all has been added, allow the mixture to stand for three days. Then pulverize and examine it. Test 1 gram of it for soluble phosphates as in experiment 106. Also test the solubility of floats in water. 109. Differences in volatility of ammonium salts. In a test tube place 2 grams of ammonium carbonate. Note the odor. Apply heat gently and observe the result. The ammonium carbonate volatilizes readily and may again deposit on the cold walls of the test tube. Try the same experiment with ammonium sulphate. This is much less volatile. In poorly ventilated barns, deposits of ammonium carbonate are frequently found, especially in horse stables. Gypsum, i.e. land plaster, when applied to the manure con- verts some of the ammonium carbonate to ammonium sul- phate and thus saves it. no. Solubility of nitrogenous fertilizers. Place 10 grams each of sodium nitrate, ammonium sulphate, and dried blood on filter papers in separate funnels. Pour over each sample small portions of water, and continue to leach the samples with water until 100 cc. of filtrate have been collected in each case. Which of these fertilizers are insoluble in water? 362 CHEMISTRY AND DAILY LIFE EXPERIMENTS TO ACCOMPANY CHAPTER XVI in. Losses from manure by leaching. Take one pound of fresh manure, place it in a large funnel having a good wad of cotton in the bottom, and leach the manure by pouring on water until one or two liters of filtrate have been collected. In a box one foot square and five inches deep, containing 15 to 20 pounds of rather infertile soil, plant 25 rape seeds. In a similar box containing the same quantity of soil, but to which the dung washings will be added as needed, also plant 25 rape seeds. After the seeds have germinated, thin down to 16 plants. Keep both boxes properly watered in a green- house or near a window and note the differences in growth. This experiment gives an idea of the great losses in plant food that may occur by the leaching of manures. 112. Loss of ammonia from manure. Place some fresh horse manure in a glass bottle and stopper it. Let it stand in a warm place for 24 hours, then open the bottle and note the odor. The penetrating odor is due to ammonia. Repeat the above experiment, but over the partly compressed manure place a layer of moist dirt one inch thick. Is the odor of ammonia so noticeable ? Moist earth is a good absorbent for ammonia. 113. Plant food in urine. Test some urine (human or animal) for potassium by the flame test as in experiment 56, also for sodium as in experiment 57. In a test tube place about 8 cc. of urine ; add a few drops of nitric acid and 5 cc. of ammonium molybdate solution. Warm to 65 C., but do not boil. A yellow coloration or precipitate shows the presence of phosphorus. Why? There is very little phos- phorus in the urine of our domestic animals. Most of the phosphates leave the body in the solid excreta. PRACTICAL LABORATORY EXPERIMENTS 363 EXPERIMENTS TO ACCOMPANY CHAPTER XVII 114. Oxygen given off by leaves. Take a small bunch of young, green, active shoots of any plant, and place it in the ap- paratus as shown in Fig. 125. Set the jar in the brightest light available, and repeat the experi- ment with another jar kept in the dark. Ar- range a third similar ex- periment in the light, with water that has been boiled and then quickly cooled. As the light falls on the leaves in the un- boiled water, little bub- bles of gas will begin to appear at the tips of the leaves. As they accu- mulate they will ascend into the test tube. No such bubbles will appear from the leaves in the dark, or from those in the light that are im- mersed in boiled water. Test the gas in the test tube for oxygen with a glowing splinter. Explain the results of these three experiments. 115. Carbon dioxide taken up by the leaves. Arrange an apparatus as shown in Fig. 126. Fill the glass tube, which is from 5 to 6 feet long, with freshly gathered leaves of grass. Use only young, active leaves. Set up the tube horizontally in the brightest daylight available, and with the aspirator slowly draw air first through the tube and then through the wash bottle containing some barium hydroxide solution. FIG. 125. Oxygen is given off by leaves. 364 CHEMISTRY AND DAILY LIFE Barium hydroxide is very sensitive to carbon dioxide, be- coming milky with a small quantity of the gas. If the light is good there will be no evidence of carbon dioxide in the air that has been drawn over the leaves. Now detach the tube of leaves and draw the ordinary air through the tube. Milk- FIG. 126. Green leaves take carbon dioxide from the air in the sunlight. iness will soon appear. This experiment shows that green leaves take carbon dioxide from the air in the sunlight. 116. Germination of seeds. Take about 20 conveniently large seeds, such as beans or corn, and soak them in water for a few hours. Then plant them in moist sawdust, cover- ing the seeds about half an inch deep. Keep them in a warm place for a day or two in order to start germination. Exam- ine the seeds from time to time and note their progress. How long does it take for the seeds to grow a radicle an inch long ? When does the plumule appear ? PRACTICAL LABORATORY EXPERIMENTS 365 117. Influence of poisons and mutilation on germina- tion. Soak half a dozen beans for a few hours and then mutilate some of them in various ways before planting them in moist sawdust. Cut into the embryo of one of the beans ; cut bits off from the embryo of another; poison a third by touching it with a trace of carbolic acid ; treat another with a drop of boiling water ; etc. Then set them out in the moist sawdust to germinate. The seed will not grow if the embryo is cut or killed. The healthy embryo itself may grow when separated from the rest of the seed, provided the embryo is kept moist with a proper food solution containing sugars, proteins, and salts. 118. The germinating seed needs air. (1) Put some mustard seed into a bottle. Fill it completely with water. Then stopper it and let it stand in a warm place. The seeds may start to sprout because of the dissolved air in the water, but they will soon stop growing. (2) Grease the stopper of a wide-mouthed pint bottle. Put into the bottle one-half ounce of mustard seed and enough water to moisten the seeds thoroughly. Stopper the bottle and set it in a warm place in the dark. The seeds will germinate, for there is some air in the bottle ; but they will soon stop growing further for lack of air. Open the bottle and at once insert a lighted taper. The flame is extinguished, indicating that the oxygen in the bottle has been used up. (3) Decant some of the air from the bottle just used into a clean bottle and shake it with clear limewater. The solution becomes milky, indi- cating the presence of carbon dioxide. Explain the results. 119. Carbon in plants. Take some dry plant tissue from growing plants, like beans or wheat, and place it in a porce- lain dish. Heat it strongly over a Bunsen flame, covering the greater part of the dish with a porcelain or iron cover. Thick gases will come off, and these will take fire if allowed 366 CHEMISTRY AND DAILY LIFE to come into contact with the flame. After a time, extin- guish the flame, by completely covering the basin. With- draw the burner, and allow the dish and its contents to cool. It will be seen that the interior of the dish is covered with a black soot which is carbon, set free from the plant tissue. 1 20. Nitrogen in plants. The important element nitro- gen, which is contained in the plant in combination with other elements, is lost during combustion and therefore is not in the ash of a plant. To show its presence in the plant mix a gram or two of dry plant material in a test tube with soda-lime, and then heat the whole directly in a flame; the gases which come off will contain ammonia, which can be smelled. Test the gas also with moist red litmus paper. 121. Phosphorus in seed. Crush 25 kernels of wheat in a mortar. Place the material in an iron or porcelain cru- cible and heat it strongly over the Bunsen burner. Do not put the cover on. When cool, transfer the charred mass to a small beaker. Add 10 cc. of nitric acid and 50 cc. of water, and boil for 10 minutes. Break up the charred particles with a stirring rod during the boiling. Filter, and to half of the warm filtrate add 3 cc. of ammonium molybdate solution. A yellow precipitate shows the presence of phosphorus. Why ? 122. Calcium in clover. Cut up some clover hay with a pair of shears. Place 25 grams of it in an iron crucible and ignite it over the burner till only ash is left. Transfer the ash to a beaker. Add 50 cc. of water and 5 cc. of nitric acid and then heat and finally filter. To about 20 cc. of the filtrate add ammonia water until the solution just turns red litmus blue. Then add 5 cc. of acetic acid, or more if necessary, until the reaction is distinctly acid to litmus. Now add 2 to 4 cc. of a saturated solution of ammonium oxalate. A white precipitate shows the presence of calcium in the ash. The precipitate is calcium oxalate, CaC 2 O4. PRACTICAL LABORATORY EXPERIMENTS 367 123. Potassium in plants. Test the ash from the clover for potassium with the flame test and blue glass as described in experiment 56. 124. Preparation of pectin. Grate a turnip and place the material in a linen bag and wash out all the soluble matter. Place the washed pulp in dilute hydrochloric acid (1 part of concentrated acid to 15 of water) and let stand 48 hours. Then squeeze out the acid liquid. Filter the liquid and to the filtrate add an equal bulk of alcohol. The precipitate formed is pectin. The pres- ence of this substance in fruits is the cause of their jellying. 125. Acids in foods. Vine- gar contains acetic acid. Taste some vinegar, and then place 25 cc. of it in a beaker. Add two drops of phenol- phthalein solution, and then carefully add a dilute solution of baking soda from a bu- rette, Fig. 127, until a red color just appears. Now taste the liquid. It is no longer acid in taste. The soda has combined with the acid to form sodium acetate, a salt which has but little taste. What gas escaped? Baking soda is often used to neutralize excessive acidity in cookery, especially when acid fruits, such as apples, plums, or lemons, are em- ployed. FIG. 127. Neutralizing an acid. 368 CHEMISTRY AND DAILY LIFE EXPERIMENTS TO ACCOMPANY CHAPTER XVIII 126. Digestion of proteins. Get a hog's stomach at a slaughterhouse. Wash the stomach thoroughly with water. Remove the mucous membrane. Grind this membrane in a sausage grinder and place it in 500 cc. of 0.2 per cent solu- tion of hydrochloric acid. Add 10 cc. of chloroform to pre- vent putrefaction. After standing at ordinary temperatures for 24 hours, filter the solution. Into 100 cc. of this solution place about 10 grams of coagulated egg white made by plac- ing egg white into boiling water. Set the dish and its con- tents in a warm place, not above 100 F., for 24 hours. Examine the result. The protein has partly, if not entirely, disappeared. It has been digested by the pepsin from the stomach. 127. Action of malt on starch. In a mortar crush 30 malted barley kernels. Put this powder in a small flask, and add 15 cc. of water. After 24 hours filter off the solu- tion, and to the liquid add 2 grams of flour and 100 cc. of water. Se"t the flask away in the cupboard for 24 hours. Then filter off the solution, and test a portion of the residue for starch with iodine solution. Malt converts starch to sugars. Test some of the solution for sugar with Fehling's solution as in experiment 54. If brewer's malt cannot be had, the barley grains themselves may be germinated ; and when the sprouts are one-half to one inch long the seeds can be used in this experiment. 128. Action of saliva on starch. To obtain the saliva for this experiment, chew a small piece of pure paraffin, and collect the saliva in a beaker. Make some starch paste by suspending about 10 grams of starch in cold water and then stirring in boiling water. Take 25 cc. of this starch paste and adjust its temperature to about 104 F. Add 5 drops of PRACTICAL LABORATORY EXPERIMENTS 369 saliva and stir thoroughly. At intervals of a minute, remove a drop of the solution to a white tile and test the drop by means of a drop of iodine solution. If the blue color with iodine still forms after five minutes, add another 5 drops of saliva and stir. The turbidity of the starch solution should soon disappear, indicating the formation of soluble com- pounds which give no blue color with iodine. After further action, test the solution with Fehling's solution. A reddish brown precipitate with this reagent shows that the starch has been converted to the sugar, maltose. EXPERIMENTS TO ACCOMPANY CHAPTER XIX 129. Water in foods. Mark two porcelain crucibles with your initials, using a lead pencil. Place them in the flame and heat them to redness for five minutes. Cool them in a desiccator and weigh them. Into each crucible place 4 grams of flour. Place them in an oven at 95 to 100 C. and allow them to remain for five hours. Cool them in the desiccator for one-half hour, and weigh them. The loss in weight, divided by the weight of the sample (4 grams) and multiplied by 100 gives the per cent of water in the flour. Flour will contain from 10 to 15 per cent water. Repeat the experiment with grated turnips, or with butter. In- stead of porcelain crucibles ordinary tinned iron salve box covers may also be used. 130. Starch and protein in bread. Make the iodine (starch) and protein tests on samples of bread. Take some of the brown crust of the bread and grind it up and extract it with water. Test this extract with Fehling's solution. If you get a test for sugar, make a similar test, using an ex- tract made from a part of the interior of the same loaf of bread. Compare the results. What are your conclusions? 2s 370 CHEMISTRY AND DAILY LIFE 131. Effect of heat on potato starch. With the point of a knife scrape lightly the surface of a raw potato and place a drop of the starchy juice upon a microscope slide. Cover this liquid with a cover glass and examine it under the com- pound microscope. Draw the starch grains. Make this examination, using a boiled potato and also a baked potato, drawing the starch grains in each case. Baking and cook- ing rupture the cells and thus it becomes easier to digest the food they contain. 132. Starch grains. With a compound microscope ex- amine the starch grains of wheat, corn, rice, and oats. Draw the grains and compare these drawings with those of the dif- ferent starch grains pictured in Figs. 86, 87, and 88. Pro- ceed as follows : Place a drop or two of the starch water made from the grain on a slide. Cover the liquid with a cover glass, and examine the material under the mi- croscope. 133. Breakfast foods. Under the microscope, examine two samples of cereal breakfast foods. Compare what you see with the appearance of the starch grains of wheat, corn, and oats, experiment 132. Tell from what grains the break- fast foods were made. 134. Eggs. Composition of eggshell. Examine a por- tion of the shell under the microscope, and note its physi- cal character. Crush and grind an eggshell. Extract it thoroughly with warm water, and then dissolve the extracted mass in dilute hydrochloric acid. Observe that a gas comes off. Hold in the gas a drop of limewater on the end of a glass rod and note how cloudy the liquid becomes. Filter the acid solution, add ammonia water until strongly alka- line, and then add 5 cc. of ammonium oxalate solution. What is the precipitate ? What is the gas that was formed ? Why must chickens be fed limestone? Try the same experiment PRACTICAL LABORATORY EXPERIMENTS 371 with a piece of limestone; and also with wood ashes in place of eggshell. 135- Egg albumen. Prick two holes in an egg and blow out the white. Shake up a portion of it with water and note that it dissolves. Then boil some of the solution. The white coagulates. White of egg is a protein, and serves to nourish the growing chick during the incubation of the egg. 136. Yeast. On a watch glass mix thoroughly a piece of yeast of the size of a grain of wheat with about 5 cc. of water. Then with a stirring rod place a drop of this solution on a microscope slide, adding a drop of very dilute methyl violet solution. Cover with a cover glass, and examine under the microscope. The living, active cells appear colorless, while the decayed and lifeless ones are stained. Yeast cells are circular or oval in shape. 137. Carbon dioxide from baking powder. Baking powders are used for the production of carbon dioxide. Place 20 grams of the dried powder in a 250-cc. flask, provided with a cork through which passes a delivery tube, having its outer end below the surface of 100 cc. of limewater placed in a narrow beaker. To preserve the limewater from contact with the air, cover with one-fourth inch of kerosene. When water is added to the baking powder, carbon dioxide gas is rapidly evolved. The gas produces a precipitate in the lime- water. What is the precipitate? 138. Alum in baking powders. Burn two grams of the baking powder in a porcelain dish. Extract the ash with boiling water and filter. Add to the filtrate enough am- monium chloride solution so that the mixture smells distinctly of ammonia. The appearance of a white, flocculent pre- cipitate, especially upon warming, indicates the presence of alum in the powder. What is the precipitate? 139. Testing baking powders for phosphates. Dissolve 372 CHEMISTRY AND DAILY LIFE one-half gram of baking powder in 5 cc. of water and add 3"cc. of nitric water. Filter, and add 3 cc. of ammonium molyb date solution and warm gently. A yellow precipitate indi- cates the presence of phosphates. 140. Fat in meat. Grind up some meat in a meat grinder. Place 20 to 30 grams in a flask and cover it with 50 cc. of gasoline. Stopper the flask loosely. Shake it occasionally and let it stand one-half hour. Pour the gasoline off into a beaker and allow it to evaporate slowly. For this purpose it may be set outdoors. Observe the fat left behind in the beaker. Do the same experiment with peanuts instead of meat, grinding them up well before pouring on the gasoline. What are the results ? 141. Action of iron compounds on tannic acid. Make an infusion of tea by placing 5 grams of tea in 100 cc. of boiling hot water and stirring well. Filter off some of the infusion, and add 5 cc. of ferrous sulphate solution (made by dissolving 1 gram of ferrous sulphate in 10 cc. of water and filtering). The solution turns dark in color ; and if concentrated enough it becomes black. Try the same experiment with an infusion of hemlock or oak bark. This is the way common ink is made, only nutgalls are used instead of tea, or the bark of trees. All of these plant materials contain tannin. EXPERIMENTS TO ACCOMPANY CHAPTER XX 142. Microscopic examination of milk. Place a drop of milk on a microscope slide. Cover with a cover glass and examine the milk for fat globules and for impurities, such as hair, dirt, etc. Make drawings. Examine a drop of skimmed milk. How do the two differ ? 143. Acidity of milk. Test fresh milk with blue litmus paper. Let some of the milk sour and test this with blue PRACTICAL LABORATORY EXPERIMENTS 373 litmus paper. The acid formed is lactic acid, produced from the sugar in the milk by bacteria. 144. Formation of lactic acid. Place 5 grams of milk sugar in 100 cc. of water, add 5 cc. of skimmed milk and 2 or 3 crystals of sodium phosphate. Add 2 or 3 cc. of blue litmus solution. Leave the flask uncorked in the cupboard for 24 hours. Has the solution turned red? If it has, lactic acid (the acid in sour milk) has been formed. Taste the solution. 145. Babcock test for fat in milk. Measure with the pipette into the test bottle 17.6 cc. of milk. The sample should be carefully taken and well mixed. Add 17.5 cc. of commercial sulphuric acid (sp. gr. 1.83) from the cylinder; mix the acid and milk by rotating the bottle. Then place the test bottle in the centrifugal machine and whirl 5 minutes at a rate of 800 to 1200 revolutions per minute. Add suffi- cient hot water to the test bottle to bring the contents up to the shoulder, and whirl two minutes. Then fill the bottle with water to near the top of the graduations and whirl again for two minutes, and then read the fat. Read from the ex- treme lowest to the highest point. Each large division, as from one to two, represents a whole per cent. Each small division represents 0.2 of a per cent. Test a Holstein milk and a Jersey milk. Test the milk from the cows on your farm. 146. Formalin in milk. The instructor should place some of the preservative in milk and have the pupils detect it. To 10 cc. of milk add 10 cc. of a solution of concentrated hy- drochloric acid containing per liter 2 cc. of 10 per cent ferric chloride solution. Pour the mixture in a teacup and heat it slowly to boiling over a flame, giving the cup a slight rotary motion. If formalin is present, a violet coloration will appear. 147. Test for oleomargarine. Foam test for purity of 374 CHEMISTRY AND DAILY LIFE butter. Heat about 3 grams of the sample in a large iron spoon over a low flame, stirring constantly with a splinter. Genuine butter will boil quietly, with the production of con- siderable froth or foam. Oleomargarine and renovated butter will sputter and make much more noise. Always compare by making a simultaneous test with genuine butter. 148. Waterhouse test. Into a small beaker pour 50 cc. of sweet milk, heat nearly to boiling, and add 5 to 10 grams of butter or oleomargarine. Stir with a glass rod until the fat is melted. Then place the beaker in cold water and stir the milk until the temperature falls sufficiently for the fat to congeal. At this point the fat, if the sample is oleomargarine, can be collected into one lump by means of a rod. If it is butter, it will granulate and cannot be thus collected. EXPERIMENTS TO ACCOMPANY CHAPTER XXI 149. Purity of Paris green. Take about one gram of Paris green. Put it in a beaker and add 25 cc. of ammonia water. Stir and let it stand for five minutes. If the " green " is pure, it will form a clear, dark blue solution, and leave no solid residue. 150. Preparation of lime-sulphur. Slake 25 grams of lime in 200 cc. of water. Add 50 grams of flowers of sulphur. Shake the mixture vigorously, and heat it in a steam bath with frequent shaking. Finally place it over a wire gauze or a free flame and boil for half an hour. The mixture must be watched and occasionally stirred in order to avoid violent bumping. After boiling, allow all to stand and settle. Note the color of the liquid and the character of the sediment. 151. Preparation of Bordeaux mixture. Dissolve 0.5 pound of copper sulphate in 2 pounds of hot water. Also slake 0.3 to 0.4 pound of lime in 2 pounds of water and strain PRACTICAL LABORATORY EXPERIMENTS 375 the liquid through a cheesecloth into a pail so as to remove the coarse material. The solution of copper sulphate is then poured into the pail and all is well stirred. In the preparation of Bordeaux mixture it is the aim to use just enough lime to combine with all of the copper sulphate. Test the mixture with litmus paper. If it is acid, enough more lime should be added to turn the litmus blue, The ferrocyanide test can also be performed as follows : Filter some of the finished mixture, and add a few drops of potassium ferrocyanide. If a reddish brown precipitate or color appears, more lime should be added. What is the reddish brown precipitate ? 152. Destruction of germinating power by formalin. Place about 6 or 8 kernels of different grains, such as corn, barley, and wheat, in a solution of formalin of 20 per cent strength. Leave them there for half an hour. Remove the seeds, wash with water, and place them in moist sawdust for germination. Do they germinate? What care must be practiced in treating seeds with such poisons? LISTS OF APPARATUS AND CHEMICALS NECESSARY FOR THE EXPERIMENTS DESCRIBED IN THIS CHAPTER GENERAL LABORATORY APPARATUS 1 Babcock tester. 1 Microscope. 1 Inexpensive balance sensitive to 1 eg., and set of weights, 50 gr. to 1 eg. 1 Drying oven (copper). 1 Water bath (6 holes). 2 Mortars (4 in.), preferably of porcelain, though glass mortars will do. 2 Thermometers (0-120). 4 Burettes, 50 cc., and holders. 2 Condensers (Liebig). 376 CHEMISTRY AND DAILY LIFE 1 Gross corks, sizes 7, 8, 9, 10, and 12. 1 Cork borer, set of six. 2 500-cc. cylinders. 6 Cobalt glasses 10 cm. square. 1 Magnet (horseshoe). 4 Pneumatic troughs, galvanized iron, 1 ft. long, 8 in. high, and 8 in. wide, or can use 1-gal. stoneware milk pans. 3 Platinum wires, 4 in. long, about No. 26 B. & S. guage. 1 Blast lamp, bellows and rubber tubes. 10 Ib. Glass tubing, sizes up to f in. int. diam. 4 Retorts (glass) 500 cc. 1 Dozen cylinders 100 cc. Tacks, nails, brads, carpenters' tools, etc., which can always be obtained from local dealers, have not been mentioned here. With the aid of carpenters' and tinners' tools many useful pieces of ap- paratus can be constructed by the students. Often the local tinner will make a piece of apparatus that will serve as well as one secured elsewhere at higher cost. LIST OF APPARATUS USED BY EACH STUDENT 1 Bunsen burner and tubing, or alcohol lamp. 2 Stirring rods. 4 Watch glasses (diam. 3 in.). 2 Erlenmeyer flasks (300 cc.). 25 Filter papers (11 cm.). 1 Box matches. 1 Wire gauze (4 in. square). 2 Crucibles (porcelain, diam. 3.5 cm.) 1 Test tube brush. 2 Funnels (diam. 6.5 cm.). 1 Sand bath (4 in. diam.). 1 Test tube stand. 1 Wooden stand for funnels. 12 Test tubes 12 cm. long, diam. about 1.7 cm. 6 Beakers (nest 100 to 600 cc.). 1 Casserole (200 cc.). 2 Evaporating dishes (diam. 7.5 cm.). 1 Crucible tongs (9 inches). PRACTICAL LABORATORY EXPERIMENTS 377 1 Test tube clamp. 1 Water bath (copper). 2 Aluminum dishes (2 in. diam.). 2 Feet of glass tubing (soft), ext. diam. 6 mm. 1 File (15 cm. long), triangular. 1 100-cc. cylinder. 1 Bottle each of blue and red litmus paper. 1 Ring stand and rings. 1 Deflagrating spoon. 4 Bottles (wide mouth) 250 cc. Rubber tubing 2 ft. (int. diam. 5 mm.). 2 Flasks 500 cc. 1 Thistle tube. 1 Desiccator. 1 Iron crucible 50 cc. CHEMICALS FOR TEN STUDENTS GRAMS Acid, Hydrochloric, sp. gr. 1.2 . . . . 1500 Nitric, sp. gr. 1.4 . . . . 1000 Sulphuric, sp. gr. 1.84 . . . . 2000 Acetic (50%) 200 Alcohol 100 Alum (potassium) . 25 Aluminum chloride 25 Ammonium carbonate 50 Ammonium chloride 250 Ammonium hydroxide, sp. gr. 0.9 . . . , 1500 Ammonium oxalate 25 Ammonium nitrate 100 Ammonium sulphate . . . . . .100 Ammonium molybdate 50 Arsenic trioxide 100 Barium chloride 20 Barium hydrate 20 Bone black 100 Bone ash . . : 100 Borax 50 Calcium chloride (granular) . . . . . 200 378 CHEMISTRY AND DAILY LIFE GRAMS Calcium sulphate (gypsum) . . . . . 50 Calcium carbonate (marble) 50 Carbon bisulphide 30 Citric acid . 25 Charcoal, ten pieces 8 cm. X 4 cm. Cochineal 5 Cobalt nitrate 10 Copper (turnings or scrap) . . . . 250 Copper oxide (black) 50 Copper sulphate 100 Formalin 50 Fuchsine ......;. 25 Glucose (sirup) 1000 Iron chloride . . . . . . . 30 Iron powder 25 Iron sulphate (ferrous) 75 Iron sulphide 250 Iodine 10 Indigo 15 Lime 50 Logwood 10 Lead acetate . 25 Lead (cuttings from lead pipe) . . . .100 Lead nitrate 10 Litmus 10 Magnesium carbonate 50 Magnesium sulphate 20 Magnesium wire or ribbon 10 Manganese dioxide 1000 Methyl violet ....... 5 Mercury 100 Mercuric chloride 20 Paraffin 10 Phenolphthalein 5 Phosphorus, yellow 10 Potassium carbonate 50 Potassium chlorate 250 Potassium sulphate 20 Potassium dichromate 25 PRACTICAL LABORATORY EXPERIMENTS 379 GRAMS Potassium hydroxide 50 Potassium iodide . 50 Potassium nitrate . 25 Potassium permanganate 10 Potassium ferrocyanide 25 Rochelle salts 20 Silver nitrate . 10 Sodium (metallic) 5 Sodium bicarbonate 20 Sodium carbonate 100 Soda-lime . .... .100 Sodium hydrogen phosphate . . . . . 20 Sodium hydroxide ....... 400 Sodium nitrate 200 Sodium silicate (water glass) . . . . . 100 Sodium sulphate 10 Strontium chloride 10 Sulphur 200 Tannic acid . . . . . . . . 20 Tin 30 Zinc (granulated) . . . . . . . 1000 Zinc sulphate (crystals) 100 This list does not include common materials that are always easily obtained, as sugar, salt, lard, tallow, clay, starch, cotton, iron, wire, gasoline, benzine, chloroform, bleaching powder (chloride of lime), tartar emetic, turpentine, etc. For the convenience of teachers the names of a number of firms from whom apparatus and chemicals may be purchased are here given: E. H. Sargent & Co., Chicago, 111.; Eimer and Amend, New York City ; Bausch and Lomb, Rochester, N. Y. ; Arthur H. Thomas Co., Philadelphia, Pa. ; Scientific Materials Co., Pittsburg, Pa. ; Henry Heil & Co., St. Louis, Mo. ; Denver Fire Clay Co., Denver, Col. ; Kny-Scheerer Co., New York City. INDEX Acetic aldehyde, 99. Acetylene, 93. Acid, acetic, 41, 98. benzoic, 104. boric, 42, 74. butyric, 105. carbolic, 103. carbonic, 89. citric, 105. definition of, 48. formic, 98. hippuric, 105. hydriodic, 57. hydrobromic, 56. hydrochloric, 53, 336. hydrofluoric, 52. lactic, 41, 104. malic, 41, 104. muriatic, 53. oxalic, 41, 104. phosphate, 208, 220, 361. phosphoric, 70. Prussic, 123. pyrogallic, 104. salicylic, 101. silicic, 42, 78. sulphuric, 61. sulphurous, 63. tartaric, 41, 104. valeric, 105. Acids, 41. in fruits, 335. in plants, 252. in silage, 335. Acidulated rock, 220. Air, carbon dioxide in, 30. composition of, 30. in the lungs, 262. moisture in, 31. nitric acid in, 35. Ajax, 276. Alabaster, 130. Albumen, 253. in milk, 294. Albuminoids, 114. Albumins, 114. Alcohol, absolute, 112. amyl, 102. denatured, 101. grain, 98. wood, 98. Alfalfa, 287. Alkalies, 42. metals of, 118. Alkaline earths, test for, 347. Alkaloids, 115. Allotropic form of phosphorus, Allspice, 283. Alum, 274. ammonium, 143. baking powders, 144. potassium, 143. sodium, 143. testing for, 349. Alumina, 141. Aluminum, 140. bronze, 142. metal, 142. paint, 142. sulphate, 143. Amalgams, 153. Amides, 253. Ammonia, aqua, 38. composition of, 37. from coal, 37. from manure, 362. preparation of, 334. properties of, 38. water, 38. Ammonium bicarbonate, 126. chloride, 39. molybdate, 157. salts in soil, 39, 361. sulphate, 39, 216. 381 382 INDEX Amyl acetate, 101. alcohol, 102. Amylopsin, 262. Analysis, 48. Aniline, 104. dyes, 104. Animal foods, 284. heat, 258. Animals, 255. action in soil, 196. carbohydrate in, 257. fat in, 257. mineral matter in, 255, 256. protein in, 256. water in, 255. work by, 257. Anthracite, 84. Antidote, for arsenic, 71. Antimony, 71. Antiseptic, defined, 20. Ants, as soil formers, 197. Apparatus, list of, 375, 376. Apparatus dealers, 379. Aqua regia, 147. Aristotle, 2. Armor plate, 160. Arsenic, 71. antidote for, 71. white, 309. Arsenious oxide, 71. Asbestos, 79, 136. Ash in plants, 12, 329. Asphalt, 177. solubility of, 354. Atropa belladonna, 115. Atropine, 115. B Babbitt metal, 71, 155. Babcock, portrait, 305. Babcock test, 303, 373. Bacon, smoking, 104. Bacteria in food, 273. in fruit, 279, 280. Bagasse, 111. Baking, 271. Baking powder, 274. alum, 144, 275, 371. carbon dioxide in, 371. phosphate, 275, 371. tartrate, 275. Baking soda, 127, 274. Balanced ration, 265. Banana oil, 101. Barium nitrate, 129. sulphate, 129, 347. Barley, 270. Barytes, 129. Base, definition of, 49. Basic lead acetate, 156. Basic slag, 218, 221. Bate liquor, 185. Batteries, blue cup, 158. primary, 158. Battery carbons, 158. Baume hydrometer, 312. Bauxite, 140. Beers, 101. Beet, leaves for silage, 288. roots, 278. sugar, 109, 249. Bell metal, 152. Bengal lights, 129. Benzene, 103. Benzine, 91. Bessemer process, 165. Bioses, 108. Bismuth, 72. subnitrate, 72. Blackjack, 64, 157. Black varnish, 177. Blast furnace, 161, 162. Blaugas, 93. Bleaching powder, 55, 317, 33' Bleeding, stoppage of, 167. Blood charcoal, 81. Blood poisoning, 74. Blood stains, 356. Blue glass, 161. Blue print, paper, 167. making of, 352. Bluing, 354. Bone, ground, 218. steamed, 219. Bone black, 81. Bones, 218. phosphorus in, 360. Boracic acid, 74. Borax, 128. occurrence of, 74. uses of, 75. Bordeaux, acid test for, 316. mixture, 152, 314. preparation of, 374. Boric acid, 74, 340. INDEX 383 Boron, 74. Bottom lands, 195. Boussingault, portrait, 201. Bowlders, 193. Bran, 271. Brass, 152, 159. copper in, 350. Bread, 270. making, 274. protein in, 369. starch in, 369. Breakfast foods, 276, 370. Brewer's grains, 276. Brewing, 276. Bricks, 145. Brimstone, 59. Britannia metal, 154. Bromine, preparation of, 55. properties of, 56. vapors, 337. Bronzes, 152. Brucine, 115. Buckskin, 185. Bullets, 72. Bunsen burner, 85. Butter, 299. composition of, 300. fertility in, 230. sweet cream, 299. Cabbage, 278. Cadaverine, 114. Calcined soda, 126. Calcium, acid phosphate, 274. carbide, 93. carbonate, 130. chloride, 130. nitrate, 132. phosphate, 66, 132. silicate, 133. Calomel, 153. Camembert cheese, 302. Cane sugar, 108, 249. inversion of, 109. Cunning of fruit, 279. Caoutchouc, 187. Capillarity, 209. Capillary water, 210, 358. Caramel, 110, 272. Carbohydrate, 105, 247. digestion of, 262. Carboleum, 318. Carbolic acid, 103, 318. Carbon, 81. absorbing power of, 341. bisulphide, 64, 314. cycle, 88. dioxide, preparation of, 341. dioxide in air, 30, 333. monoxide, 86. preparation of, 341. reducing power of, 341. tetrachloride, 95. Carbonated waters, 88. Carbonate of potassium, 120. Carbonates, 89. Carboneum, 95. Carnallite, 119. Carrots, 278. Casein, 294, 304. Cassiterite, 153. Cast iron, 163. Caustic soda, 125. Cells, 238. Celluloid, 107. Cellulose, 105, 250. nitrates of, 106. Cement, for floors, 131. natural, 134. Portland, 133. Centrifuge, 111. Cereals, 270. composition of, 270. Cerium, 146. Chalk, 130. Chamois skin, 185. Charcoal, 81. Cheese, 300. composition of, 302. Chemical change, 4, 326, 327. Chemical equations, 45. Chemicals, list of, 377. Chili saltpeter, 36, 121, 217. Chloride of lime, 55, 317. Chlorine, 53. Chlorine water, 55. Chloroform, 95. Chlorophyl, 168, 245. Chrome green, 156, 174. steel, 156. tanning, 185. yellow, 156, 174. Chromium, 156. Churning, 299. 384 INDEX Cider, hard, 284. Cinchonine, 115. Cinnabar, 153, 174. Cinnamon, 283. Cisterns, 16. Clay, 79, 144, 192, 198. puddling of, 357. settling of, 359. Cloth, dyeing of, 354. Clouds, 10. Clover, 287. calcium in, 366. Coal, 83. Coal gas, 84. Coal tar dyes, 104. Cobalt, 160. bead, 351. Cocaine, 115. Cocoa, 282. Codeine, 115. Coefficient of digestibility, 263. Coffee, 282. stains, 356. Coins, silver, 48. Coke, 81. Colemanite, 75. Collodion, 36, 106. Collostrum, 295. Common salt as fertilizer, 223. Compound, denned, 6. Compounds, 326. Concrete, 134. reenforced, 134. Condensed milk, 302. Contact poisons, 308, 311. Cooking, 271. effect of, 272. Copal, 176. Copper, 150. acetate, 100. alloys of, 148, 150. cleaning of, 152. in coin, 349. plating, 152. properties of, 349. Copperas, 166. Corks, fitting of, 322. Corn, 270. Corn, popping of, 273. "Corn flakes," 277. Corn starch, 275. Corn stover, 288. Correcting soil acidity, 357. Corrosive sublimate, 153, 316. Cottage cheese, 301. Cotton, 180. burning of, 353. Cotton seed meal, 217. Cotton seed oil, 102. Cream, 297. Cream of tartar, 104, 274. Cream separators, 298. Creosol, 104, 318. Creosote, 98, 104, 318. Crude fiber, 247. Cryolite, 141. Cyanides, 313, 314. Davy, portrait, 124. Definite proportions, law of, 44. Denatured alcohol, 101. Depilation, 183. Developer in photography, 149. Dew, 10, 32. Dextrine, 107, 249. Dextrose, 108. Dialysis, 78. Diamond, 82. Diastase, 112, 276. Digestibility of feeds, 263. Digestion, 258. of proteins, 368. Dioxygen, 28. Distillation, destructive, 38. dry, 38. Distilled water, 10, 328. Distiller's grains, 276. Dolomite, 89, 136. Double decomposition, 48. Drainage, 210. Dried blood, 217. Dried fruit, 281. Driers, 172. Dry batteries, 158. Dry farming, 209. Dutch metal, 152. Dyestuffs, 104, 182. Dynamite, 36, 80, 106. Dysentery, remedy for, 72. E Earthenware, 145. Earthworms, action of, 196. INDEX 385 Ebonite, 189. Egg albumen, 371. Eggs, 286. composition of, 370. Elastin, 114. Elements, ancient idea of, 2. chemical, list of, 3. classification of, 6. defined, 2. essential, in plants, 119. Emery, 140. Enzymes, 259. Epsom salts, 66, 137. Equivalents, hydrogen, 45. Esters, 101. Etching glass, 52. Ether, 103. Ether extract, 251. Face powder, 137. Fat, digestion of, 262. in meat, 372. Fats, 250, 251, 345. Fattening animals, 266. Feeding standards, 263. Fehling's solution, 108, 346. Feldspars, 79, 119. Fermentation, alcoholic, 342. Ferments, unorganized, 114. Ferric chloride, 116, 167. oxide, 116. Ferrous chloride, 116. oxide, 116. sulphate, 116. Fertility, factors of, 202. Fertilizer, phosphatic, 218. Fertilizers, action of, 360. analysis of, 226, 227. commercial, 214. elementary system, 227. experimenting on, 224. home mixing of, 224. indirect, 222. laws on, 226. mixed, 223. nitrogenous, 216. selection of, 224. special, 224. Filtering, 323, 324. Fire bricks, 146. Fire clay, 146. 2c Fires, extinguishing, 96. Fish, as fertilizer, 217. as food, 286. smoking, 104. Fixation of nitrogen, 207. Flashlight powder, 136. Flaxseed, 172. Floats, 207, 219. Flour, 107, 271. Fluorine, 52. Fly removers, 318. Fogs, 10. Foods, acids in, 367. animal, 269. human, 269. water in, 369. Foodstuffs, classes of, 246. Fool's gold, 64. Forage, 287. " Force," 276. Formaldehyde, 98. Formalin, 98, 316. action on seeds, 375. in milk, 373. preparation of, 351. Formates, 98. Fractional distillation, 91. Fructose, 108, 250. Fruit stains, 355. Fruit sugar, 108. Fruits, 278. Fundamental considerations, 1. Fungicides, 314. Furs, 186. Galenite, 64, 154. Galvanized iron, 157. Gasoline, 91. Gelatin, 286. German silver, 152, 160. Germination, 239, 240, 364. Germol, 318. Germs in water, 17. Gin, 112. Glaciers, 193. Glass, 134, 135. quartz, 77. Glass tubing, bending of, 322. cutting of, 321. Glauber's salt, 128. Glazes, 145. 386 INDEX Globulin, 253. Glucose, 108, 249. Glue, 286. Gluten, 107. Gluten feed, 276. Glycerine, 92, 101. oleate, 102. palmitate, 102. stearate, 102. Glycogen, 257. Gold, 147. copper alloys of, 48. Granitic rocks, 79. "Grape-nuts," 276. Grape sugar, 108, 249. Graphite, 82. Grass stains, 256. Gravitational water, 210. Grease spots, 251, 343. Green manuring, 236. Green vitriol, 166. Growing animals, 266. Guajacol, 104. Guano, 218, 219. Guncotton, 36, 106. Gun metal, 152. Gunpowder, black, 121. smokeless, 106. Gutta-percha, 189. Gypsum, 64, 130. as fertilizer, 216. H Hail, 32. Hair, as fertilizer, 217. Halogens, 52. Ham, smoking, 104. Hammer black, 166. Hard soap, preparation, 342. Hay, 287. Heat as disinfectant, 319. Heavy spar, 219. Hellebore, 310. Hematite, 161. Hemoglobin, 168. Hippuric acid, 206. Hoof meal, 217. Hornblende, 79, 136. Horn meal, 217. Human foods, 269. Humus, 199, 212. Hunyad spring, 137. Hydrocarbons, 91. Hydrochinone, 104. Hydrogen, dioxide, 28. equivalent, 45. occurrence, 22. peroxide, 28. preparation of, 330. properties of, 22. sulphide, 63, 339. Hydrocyanic acid gas, 313. Hydroscopic water, 210. Hypo, 128. Hyposulphite of soda, 128. Ice cream, 302. Indian corn, 275. Infusorial earth, 80. Ink, 167, 352. Ink stains, 355. Insecticides, gaseous, 313. Insect powder, 311. Invertase, 110. Iodine, 56. action on starch, 57. vapors, 337. [odoform, 95. Iron, change to ferric, 352. compounds with tannic acid, 372. metal, 161. pyrites, 64. rust, removal of, 354. rusting of, 166. sulphate, 352. ridium, 147. singlass, 286. soprene, 187. apan driers, 173. elly making, 250. unket tablets, 294. K finite, 119, 222. Calsomine, 177. taolin, 79, 144. Ceratin, 114. Cerosene, 91. emulsion, 312. Cohlrabi, 278, INDEX 387 Lactic acid, 373. Lactometer, 340. Lactose, 108, 113. Lampblack, 81. Land plaster, 131, 223. Lard, hogs', 120. Latex, 186. Lead, 154. acetate, 100. action of water on, 351. alloys of, 154. arsenate, 71, 156, 310. chr6mate, 351. colic, 156. salts of, 350. storage battery, 155. sugar of, 156. Leaf, action of, 242, 246. Leather, bating, 185. box calf, 186. gelatin in, 354. kid, 185. making of, 183. Morocco, 184. Russia, 185. soft, 184. sole, 185. split, 185. tawing process, 185. waste, 217. Leaves, carbon dioxide from, 363. evaporation from, 245. oxygen from, 363. LeBlanc soda process, 125. Lecithin, 67. Legumes, use of, 236. Levulose, 108, 109. Lice exterminator, 318. Liebig, portrait, frontispiece. Lime, 130. air slaked, 223. as disinfectant, 317. deposits of, 329. quality of, 348. value of, 222. Limestone, 89, 130. as fertilizer, 216. ground, 223. weathering of, 357. Lime sulphur, 311, 374. Limonite, 161. Linen, 182. Linseed oil, 171, 172, 353. Lipase, 262. Liquids, heating of, 325. pouring of, 323. Litharge, 155. Lithium, 118, 123. Lithophone, 175. Loam, 198. Loess, 195. London purple, 310. Lubricating oils, 91. Lunar caustic, 148. Lysol, 318. M Magnalium, 142. Magnesium, 136. burning of, 348. carbonate, 137. chloride, 137. oxide, 137. Magnetite, 161. Malleable iron, 164. Maltose, 108, 112, 250. Manganese, bronze, 159. dioxide, 159. metal, 159. salts, 160. Manure, ammonia in, 233. cold, 229. composition of, 230. effect in field, 234, 235. fermentation of, 233. hot, 230. increasing value of, 231. kinds of, 229. leaching of, 362. liquid, 232. losses of, 232, 233. reenforcing of, 231. saving of, 233, 234. spreader, 235. storing of, 234. value of, 229. Marble, 89, 130. Marsh gas, 93. Mastication, 259. Matches, 69. safety, 69. Swedish, 69. Meal, 271. 388 INDEX Meat, 284. Meat scraps, 217. Medium ratio, 264. Meerschaum, 79, 136. Mercuric chloride, 316. Mercury, 152. antidote for, 350. fulminating, 106. preparation of, 350. Metalloids, 6. Metals, 6. base forming, 49. Methane, 93. Methyl alcohol, 97. salicylate, 101. Mica, 79. Milch cows, needs of, 267. Milk, 291. acidity of, 372. antiseptics in, 296, 297. ash in, 295. composition of, 292. deep setting, 297. digestion of, 261. fat in, 293. fertility in, 230. germs in, 295, 296. globules, 294. lactic acid in, 295. powders, 302. preservatives, 297. secretion of, 291. shallow setting, 297. souring of, 294. sugar, 108, 294. yield, 291. Mill concentrates, 270. Mixtures, 326. Moisture in air, 31. Molasses, 111. Molds in fruit, 281. Molybdenum, 157. Monoses, 108. Mordants, 167, 182. Morphine, 115. Mortar, 90. Mosaic gold, 154. Moth balls, 318. Muck, 198. Mulch, 209. Muriate of potash, 119, 222. Muriatic acid, 53. Mutton tallow, 102. N Naphtha, 91. Naphthalene, 318. Narcotine, 115. Narrow ratio, 264. Negative, photographic, 150. Nernst lamp, 146. Neutralization, 42. Neutral oil, 300. Nickel, ammonium sulphate, 160. coins, 160. metal, 160. plating, 160. Nicotine, 115. Nitrate of soda, 217. Nitrates, 36, 205. Nitric acid, 35, 36, 334. Nitrification, 205. Nitrites, properties of, 37. Nitrobenzene, 104. Nitrogen, 32, 34, 333. cycle, 206. fertilizers, 361. free extract, 247. in soil, 356. tests for, 344. Nitroglycerine, 106. Nitrous acid, 37. Non-metals, 6, 50. Nucleoproteins, 114. Nut butter, 282. Nutmeg, 283. Nutritive ratio, 264. Nuts, as food, 272. "Nuttolene," 282. Nux vomica, 115. O Oatmeal, 277. Oats, 270. Oil gases, 93. Oil of mirbane, 104. Oil of vitriol, 61. Oils, 250. drying, 171. non-drying, 171. solubility of, 353. Oleomargarine, 300, 373. Oleo oil, 300. Olive oil, 102. Onion, 278. IXDKX 389 Open hearth process, 165. Opium, 115. Organic compounds, 344. Osmium, 147. Oxides, acid forming, 26. alkali forming, 26. Oxygen, 24, 331. Oxyhydrogen gas, 327. Oysters, 284. Ozone, 27, 333. Paint, 171. enamel, 176. stains, 355. Palladium, 147. Palm oil, 102. Paper, 105. Paraffin, 91. Paris green, 71, 309. test for, 374. use of, 310. Pasteur filters, 18. Pasteurization of milk, 296. Peanuts, 277. Pearlash, 121. Pearl barley, 276. Peat, formation of, 198. Peat bogs, 197. Pectins, 250, 367. Pepper, 283. Pepsin, 114, 260. Peptones, 114. Permanent white, 129. Peruvian bark, 115. Petroleum, 91. Petroleum ether, 91. Pewter, 154. Phenol, 103. Phenolates, 103. Phosphate of lime, 218, 220. Phosphate rock, 66, 132, 218. Phosphates, 339. Phosphor bronze, 152. Phosphorus, cycle, 68. in fertilizers, 70. in soils, 207, 356. occurrence of, 66. red, 69. yellow, 69. Photographic plate, 149. Photography, 150. Photography, developers in, 104. Physical change, 5, 326. Pig iron, 163. Pigments, 174. Pink salt, 154. Pintsch gas, 93. Pitchblende, 138. Plant, stem of, 244. Plant food, 244. in crops, 202. in urine, 362. Planting of seeds, 358. Plant life, 238. Plants, as soil formers, 197. carbon in, 365. nitrogen in, 366. potassium in, 367. Plaster of Paris, 131, 348. Platinum, 146, 147. Plumbago, 82. Poisons for pests, 307. Poppy seed oil, 171. Porcelain, 144, 145. Portland cement, 133. Positive, photographic, 150. Potash, 118. caustic, 120. fertilizers, 221, 360. red prussiate of, 167. Potassium, 118. bromide, 122. chlorate, 122. chloride, 121. cyanide, 122. hydroxide, 347. in soils, 208. iodide, 122. metallic, 123. nitrate, 121. permanganate, 160. phosphate, 122. platinic chloride, 147. silicate, 122. test for, 347. Potato, 277. Potato scab, treatment of, 98. Potato starch, 248. Producer gas, 87. Proteids, 114. Protein, coagulation of, 349. digestion of, 262. preparation of, 345. test for, 344. 390 INDEX Proteins, 114, 252. digestibility of, 272. of rye, 270. of wheat, 270. Protoplasm, 238. Prussian blue, 174. Prussic acid, 123. Ptomaines, 114. Ptyalin, 259. Puddling, 210. Puddling process, 164. Pulses, 277. Putrescine, 114. Putty, 130. Pyrethrum, 310. Pyrite, 166. Pyrogallic acid, 104. Pyrogallol, 104. Pyroligneous acid, 98. Q Quartz, 79. Quartz glass, 77. Quicksilver, 152. Quinine, 115. R Radish, 278. Radium, 138. Rain, 10. Rare earths, 146. Reaction of soil, 357. Red ocher, 161, 174. Rennet, 300. ' Rennin, 260. Renovated butter, 300. Reverted phosphate, 22. Rhigolene, 91. Rhodium, 147. Rice, 272, 276. Rochelle salt, 108. Rock, weathering of, 192. Rock candy, 109. Rock phosphate, 207. Root, action of, 242. Root hairs, 242. Root nodules, 206. Roots, 277. Roquefort cheese, 302. Rosin, 173. Rubber, 186. Rubber, black, 188. blue, 188. coats, 188. hard, 189. Para, 186. plantations, 186. preparations, 187. red, 188. solubility of, 354. substitutes, 189. synthetic, 189. vulcanized, 188. white, 188. Rubies, 140. Rum, 101, 112. Rutabaga, 278. Ruthenium, 147. S Saleratus, 127, 274. Sal soda, 125. Salt, common, 123. definition of, 49. preparation of, 336. Saltpeter, 121. Salts, examples of, 50. Epsom, 66. in soil, 358. San Jose scale, 311. Sand, 192, 198. Sand dunes, 195. Sand paper, 77. Sapphires, 140. Sap, movement of, 244. Science, definition of, 1. Seeds, air for, 365. behavior of, 241. mutilation of, 365. phosphorus in, 366. Serpentine, 79. Shellac varnish, 176. Ship's biscuit, 274. Shoddy, 182. Shot, 72. "Shredded wheat," 276. Siderite, 161. Silage, 252, 288. acids in, 290. Silica, 76. Silicates, insolubility of, 340. Silicon, 76. dioxide, 76. INDEX 391 Silk, 180. Silt, 198. Silver, 148. bromide, 148. chloride, 148. iodide, 149. nitrate, 148. plating, 150. sterling, 148. Skim milk, 297, 299. Skim milk cheese, 302. Slag, 163. Sleet, 32. Smalt glass, 161. Smokeless powder, 36. Snow, 32. Soap, 92. calcium, 102. hard, 102. making, 251. soft, 102. Soapstone, 79. Soda, 125. Soda water, 88. Sodium, acetate, 100. action on water, 331. benzoate, 104. bicarbonate, 274. carbonate, 125. chloride, 123. hydroxide, 125. iodate, 128. metallic, 123. nitrate, 127. silicate, 77, 128, 340. sulphate, 128. sulphite, 128. test for, 347. thiosulphate, 128. Soft coal, 84. Solder, 154. Solvay process, 126. Soil, 191. acids in, 207. classes of, 198. color of, 212. composition of, 200. drainage, 359. fertility of, 202. formation, 192. liming, 207. sedentary, 194. temperature of, 212, 359. Soil, transported, 194. water, 209. Spices, 283. Spirits of hartshorn, 38. Spirits of wine, 98. Springs, sulphur, 66. Stannic chloride, 154. Stannous chloride, 154. Starch, 88, 107, 247. action of saliva on, 368. effect of heat on, 370. grains, 370. in seeds, 245. paste, 107, 248. preparation of, 345. solution of, 368. test for, 346. Starter, 301. Stassfurt fertilizers, 222. salt beds, 121. Steel, 165. hardening of, 165. tempering of, 165. Steinholz, 137. Stomach, of cow, 260. of horse, 260. Stomachic poisons, 308. Storage battery, 155. Straws, 288. Strontium nitrate, 130. Strychnine, 115. Styptic cotton, 167. Sublimate, corrosive, 153. Sucrose, 108. Sugars, 108, 249. Sugar beet, 278. Sugar, in plants, 245. Sugar of lead, 156. Sugar, manufacturing of, 110. test for, 346. Sulphates, 62. Sulphate of potash, 222. Sulphites, 63. Sulphur, cycle, 66. dioxide, 59, 338. disinfectant, 317, 318. ethers, 103. flowers of, 59. in animals, 64. in plants, 65. in soils, 208. occurrence of, 59. properties of, 338. 392 INDEX Sulphur, roll, 59. springs, 66. trioxide, 61. Sulphuric acid, 338. contact process, 62. lead chamber process, 62. Nordhausen, 63. properties of, 62. preparation of, 61. Sunlight as disinfectant, 319. Superphosphate, 68, 207, 220. Swiss cheese, 302. Sylvite, 119. Symbols of elements, 43. Synthesis, 48. Talc, 137. Tannin, 167, 183. Tartaric acid, 274. Tea, 282. stains, 356. Theories, definition of, 2. Thermite, 142. Thorium, 146. Tilth, 208. Tin, foil, 153. plate, 153. stone, 153. Tobacco decoctions, 313. Trypsin, 262. Tubers, 277. Tungsten, 157. electric lamps, 157. Turnbull's blue, 167. Turnip, 278. Turpentine, 174. Type metal, 71. U Ultramarine, 146. Ultramarine blue, 174. Urea, 206. Uric acid, 206. Vacuum pans, 111. Varnish, 175. Vaseline, 91. Vegetables, fresh, 269. Verdigris, 152. Vermilion, 153, 174. Vinegar, 41, 99, 283. Vitriol, green, 166. oil of, 61. white, 159. Vulcanite, 189. W Washing soda, 126, 343. Water, chlorides in, 337. cleansers, 16. composition and uses, 8. distilled, 10. drinking, 14. electrolysis of, 8, 327. for plants, 202. gas, 86. germs in, 17. glass, 77. hard, 15, 102, 329. in animals, 13. in plants, 328. in plants and animals, 11. mineral matter in, 15. occurrence of, 328. pollution of, 18. purification of, 19. saline matters in, 9. soft, 329. spring, 14. well, 14. Waterhouse test, 374. Well water, salts in, 327. Welsbach lamp, 87. Wheat, 270. Whisky, 101, 112. White arsenic, 71, 309. White lead, 156. White vitriol, 159. Whitewash, 317. Whiting, 130. Wide ratio, 264. Wind, action of, 159. Wines, 101. Wintergreen oil, 101. Wolff-Lehmann standard, 264, 265. Wood alcohol, 98. Wood ashes, 222. Wood's metal, 72. Wool, 180. burning of, 353. Wool waste., 217. INDEX 393 Working animals, needs of, 266. Wounds, cauterizing of, 148. Wrought iron, 164. Yeast, 371. in food, 273. in fruit, 279, 280. Yellow ocher, 161, 174. Zenoleum, 318. Zinc, 157. chloride, 159. oxide, 159. oxide ointment, 159. sulphate, 159. white, 159. ; HE following pages contain advertisements of a few of the Macmillan books on kindred subjefts. Warren's Elements of Agriculture By G. F. WARREN, Professor of Farm Management and Farm Crops, New York State College of Agriculture at Cor- nell University Cloth, ismO) 4.56 pages, $/./o net Written by Professor G. F. Warren, who is in charge of the Department of Farm Management and Farm Crops in the New York State College of Agri- culture, Cornell University, an authority on questions pertaining to practical agriculture. Professor Warren is, moreover, a farmer. He grew up on a farm in the mid- dle West and is living at the present time on a farm of three hundred and eighteen acres, which he supervises in connection with his work at the Univer- sity. 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Histology, or the Minute Structure of Plants ; PART IV. The Kinds of Plants, including a Flora of 130 pages. THE MACMILLAN COMPANY Publishers 64-66 Fifth Avenue New York BOSTON CHICAGO ATLANTA DALLAS SAN FRANCISCO Shelter and Clothing: A TEXTBOOK OF THE HOUSEHOLD ARTS By HELEN KINNE, Professor of Household Arts Educa- tion, and ANNA M. COOLEY, Assistant Professor of House- hold Arts Education, Teachers College, Columbia University. In press This book and the volume, Foods and Household Management, that follows it, make up a full course in domestic matters not confined to details of cooking and sewing. The books treat fully, but with careful balance, every phase of home-making. The authors hold that Harmony will be the keynote of the home in proportion as the makers of the home regard the plan, the sanitation, the decoration of the house itself, and as they exercise economy and wisdom in the provision of food and clothing. " Home Economics stands for the utilization of the resources of modern science to improve home life," and to this end homemakers should be con- versant with modern scientific thought on matters domestic. The best schemes of heating and lighting, modern arrangements for the disposal of waste, the sanitary efficiency of tinted walls, of bare floors, of furniture built on simple lines, these are some ways in which modern science instructs the intelligent homemaker. In the selection of textiles for clothing and domestic use, a housekeeper to be efficient must be able to distinguish between fabrics of dif- ferent fibers and to choose durable weaves, she must be able to detect adultera- tion and the deceptive " finishing " processes. In buying ready-made garments she must know how to protect herself and her family from the danger of gar- ments infected by diseased operators in sweatshops. The up-to-date book on home economy treats such topics and relates them to common experience. The plan of the book is flexible. Parts may be omitted or shifted to meet the necessity or the convenience of different schools. The chapter headings in some measure disclose the breadth, the variety, and the practicability of the book : The Home. Its plan and construction ; heating, ventilating, lighting, water supply, and the disposal of waste ; decoration ; furnishing. Textiles. Materials and how they are made. Garment-making. Patterns ; cutting and making garments; embroidery. Dress. History of costume ; hygiene of clothing ; economics of dress ; care and repair of clothing ; millinery. THE MACMILLAN COMPANY 64-66 Fifth Avenue Chicago New York City Dallas Boston Atlanta San Francisco THIS BOOK IS DUE ON THE LAST DATE STAMPED BELOW AN INITIAL FINE OF 25 CENTS WILL BE ASSESSED FOR FAILURE TO RETURN THIS BOOK ON THE DATE DUE. THE PENALTY WILL INCREASE TO SO CENTS ON THE FOURTH DAY AND TO $1.OO ON THE SEVENTH DAY OVERDUE. BILK LiiAil , '' uh ~ 9 1047 iPR 11 1950 , _, -, iqSS Kftf 1 1 > JO ^PR 2 7 1956 && ^ ^ |AY ,41953 'D .-:' ' ? UG 2 1973 nrwrt-9 LD 21-100m-12, '43 (8796s) U. C. BERKELEY LIBRARIES 265529 H3 A3 UNIVERSITY OF CALIFORNIA LIBRARY