nf (Ealtfurttta ai N< GIFT OF Pacific Coast Journal of BIOLOGY r, 19J/ A TEXT-BOOK OF: CHEMISTRY if D CHEMICAL URANALTSIS EOR NURSES BY HAKOLD L. AMOSS, S.B., S.M., M.D., DE. P. H. FORMERLY CHEMIST, HYGIENIC LABORATORY, UNITED STATES PUBLIC HEALTH SERVICE ; PHYSIOLOGICAL CHEMIST, UNITED STATES BUREAU OF CHEMISTRY; INSTRUCTOR IN PHYSIOLOGICAL CHEMISTRY, GEORGE WASHINGTON UNIVERSITY MEDICAL SCHOOL, WASHINGTON, D. C. J ASSISTANT IN PREVENTIVE MEDICINE, HARVARD MEDICAL SCHOOL, BOSTON, MASS., ETC. LEA & FEBIGER PHILADELPHIA AND NEW YORK 1915 LIBRARY D ' GIFT PAG MO >, -RNAL OF NURoING TO H/Jii-.Jc DEPT.j Entered according to the Act of Congress, in the year 1915, by LEA & FEBIGER, in the Office of the Librarian of Congress. All rights reserved. Jt a' /?' PKEFACE. THE demand for a simple book on chemistry, written especially for, and adapted to, the needs of the nurse, has become more and more urgent with the institution of regular courses on chemistry in the best training schools. In the preparation of this work the author has endeavored to cover the subject briefly and clearly, and in an interesting style, so that the reader may easily absorb and assimilate the material presented. It is not to be questioned that the more chemistry a nurse knows in usable form the greater her value to the patient and to the physician. In many cases the important measure of feeding the sick is left to the nurse. In most instances her empirical knowledge of dietetics may guide her aright. Yet it is obvious that some knowledge of the chemical composition of food- stuffs and of the chemical processes of digestion and assimilation will better serve her in cases of unusual type and those presenting metabolic and digestive disturbance. If the nurse knows no chemistry, how can she be expected always to remember that starches yield sugars and are to be given to the diabetic with extreme caution? What can the term Calorie mean to her? There are other questions, too, of drug admin- 743509 IV PREFACE istration and application, of the collection and preser- vation of specimens, and of discerning observation. It is the author's impression that chemistry has not yet received the attention in undergraduate and graduate instruction of nurses that its importance merits. He believes that one of the chief reasons is that few teachers have the time to collect and put into simple form the material necessary for the nurse's foundation and understanding in chemistry. In this volume the author has striven for simplicity and a gradual logical development of the subject; and he hopes that the book will adequately fill the place in nursing education for which it was designed. H. L. A. NEW YORK, 1915. CONTENTS. CHAPTER I. THE SLAKING OF LIME .... ..... 19 CHAPTER II. WEIGHTS AND MEASURES 24 CHAPTER III. THE METALS ...... ' . 25 CHAPTER IV. MOLECULES AND ATOMS . -, . . . ..... 30 CHAPTER V. CHEMICAL PROCESSES 34 CHAPTER VI. ATOMIC WEIGHTS . 39 CHAPTER VII. OXYGEN . . . . . ..._.. . . . . . 43 CHAPTER VIII. HYDROGEN . . ' 48 CHAPTER IX. WATER vi CONTENTS CHAPTER X. HEAT ... . V* . 60 CHAPTER XI. SOLUTIONS AND PURIFICATION OF SUBSTANCES ... 64 CHAPTER XII. NATURAL WATERS CHEMICAL ACTION OF WATER . . 68 CHAPTER XIII. COMPOSITION OF WATER 72 CHAPTER XIV. HYDROGEN PEROXIDE . . . . . ... . . . . . 77 CHAPTER XV. CHLORINE ^. 81 CHAPTER XVI. BROMINE IODINE FLUORINE 89 CHAPTER XVII. SULPHUR ; . . . ..... . 95 CHAPTER XVIII. SODIUM ...... ...... . . . . . . 100 CHAPTER XIX. ACIDS AND BASES POTASSIUM . . . . . . . . 106 CHAPTER XX. PHOSPHORUS ARSENIC ANTIMONY BISMUTH 112 CONTENTS vii CHAPTER XXI. CALCIUM . . . . . . . . . 119 CHAPTER XXII. MAGNESIUM GROUP . . . . . . . . . . . 123 CHAPTER XXIII. ALUMINUM IRON MANGANESE . . ..... . 129 CHAPTER XXIV. LEAD SILVER PLATINUM . . . 133 CHAPTER XXV. CARBON . . . . . 138 CHAPTER XXVI. COMPOUNDS OF CARBON WITH HYDROGEN . . . . . 141 CHAPTER XXVII. ETHERS . . .".'.- . . . . 150 CHAPTER XXVIII. THE MARSH GAS SERIES . ......... . 153 CHAPTER XXIX. THE PARAFFINS ... . . ... .. . . - 160 CHAPTER XXX. SUGARS .... ." ; . . . . . . . . . 165 CHAPTER XXXI. POLYSACCHARIDS 173 Vlll CONTENTS CHAPTER XXXII. THE DIGESTION OF CARBOHYDRATES 180 CHAPTER XXXIII. FATS 183 CHAPTER XXXIV. BENZENE SERIES 191 CHAPTER XXXV. NITROGEN . . . . . 202 CHAPTER XXXVI. OTHER NITROGEN COMPOUNDS . . 208 CHAPTER XXXVII. PROTEINS . . ... . . . 216 CHAPTER XXXVIII. THE BLOOD 226 CHAPTER XXXIX. MILK . ... . . . 235 CHAPTER XL. THE URINE . . . 240 CHAPTER XLI. URANALYSIS . 249 TO THE INSTRUCTOR. MANY students of medicine and nurses find the particular chemistry which they need very difficult to acquire. Obviously, the manner of presentation of the subject will affect largely the results, especially in the training school where the all-important demonstration facilities are usually lacking and where there is rela- tively little time devoted to the teaching of chemistry. It is exceedingly difficult to really grasp this subject by reading or listening to lectures unaccompanied by experimental work, and the student is soon found simply attempting to memorize a large number of isolated facts. Especially is this true in the beginning, and chemistry soon becomes a hopeless muddle. Demonstrations, therefore, are urgently recommended, and whenever possible even a small number of labora- tory exercises should be arranged so that the student can handle the test-tube and chemicals. In this man- ner it is often possible to make the subject fascinating and interesting instead of drudgery. This compila- tion is necessarily limited to those subjects which the efficient nurse must know: many processes of which all educated persons should possess at least a little knowledge have been omitted. If the student acquires 2 18 TO THE INSTRUCTOR fundamental ideas of chemistry these things will come easily later in their general reading. However, it is not to be expected that the average student will digest all the matter contained herein. Enough has been left in the reduction of the original manuscript to satisfy the interested student and not too much to frighten the uninterested. An attempt is made to develop the subject from the simplest and more familiar phenomena. Since the interest of the author was first aroused in this subject by the slaking of lime this is made the foundation for the introduction of the nurse to this important science. CHEMISTRY FOR f URSES. CHAPTER I. THE SLAKING OF LIME. WHEN water is poured over quicklime, heat is gen- erated and the water soon begins to boil. If water is poured on the same amount of limestone from which the lime is made we obtain no change in temperature and no heat is generated; hence, there must be some change other than the physical contact of the lime and water in order to produce this large amount of heat. If there is such a change, what does it amount to and how is it produced? Quicklime is made by heat- ing very intensely and over a rather long period of time the stones we know as limestone rock. If we heat absolutely dry limestone in the kiln for some hours we find by reweighing it from time to time that it gradually loses weight. Continuing the heating processes we find that we arrive at a stage where it no longer loses weight, that is, we have heated it to constant weight, and no amount of heating can reduce the weight further. Curiously enough, the loss is always 44 pounds in a hundred. What does it mean then if we take several samples of pure dry 20 THE SLAKING OF LIME limestone and heat it in the kiln to constant weight ahd find that exactly 44 out of every 100 pounds is lost-? -.Can we not say, first, that the limestone probably has ' a : definite ;cqmposition that is, it is composed of certain definite materials, and second that these materials whatever they are occur in the same propor- tion each time? During this process of heating, when the limestone is being transformed into lime, the physical characters of the limestone undergo marked changes, chief among them is the change from a hard, granular substance to a smooth, friable, soft material. But what happens during the heating? It will be shown later, that this 44 pounds out of a hundred which is given off is the same gas that we give off constantly from our lungs, namely, carbon dioxide. And the driving off of this gas by heat is a chemical process and a chemical change is brought about. We could grind the original limestone into a very fine powder that would be a physical change yet on pouring water over it we would get no heat. If on the other hand we raise this powder to red heat for several hours, we have driven off a part of it though we can not see it come off and have produced a change in the internal structure of the substance (chemical change). We also witness a chemical change when water is added to the lime. Suppose that we add to our 56 pounds of lime result- ing from the heating of 100 pounds of limestone, a hundred pounds of water at room temperature (68 F.) and stir for a few minutes always the temperature of the water increases the same number THE SLAKING OF LIME 21 of degrees it makes no difference how many experi- ments you make. Then again there must be some definite and always constant change taking place when water is poured on lime. Now if we try to recover our hundred pounds of water by distillation (a physical process) we find that at the temperature of boiling water (212 F.), we can recover only 82 pounds of the 100 pounds of water which we added. Therefore 18 pounds of water must have been fixed to the quicklime in the slaking process. There must have been some chemical change also for our lime is no longer a smooth non-crystallizing substance but consists of clear crystals having definite angles. Since water or some of the constituents of water has been added to our quicklime we can call this new crystallized substance hydrated lime. This new product weighs 74 pounds and does not give off water when heated to boiling point. If now we heat our dry hydrated lime to a higher degree it begins to lose weight and water is given off until finally the whole of the 18 pounds of water is recovered and we have 56 pounds of quick- lime left, having properties identical with those before described. This is another chemical change the water was driven off by the heat leaving behind a compound with different physical properties. Suppose we add again to our 56 pounds of quicklime 100 pounds of water. When the solution has cooled and settled clear, remove a small portion in a test-tube and blow through it by means of a glass tube. The clear solution becomes turbid and a white precipitate is formed. Suppose we bubble by means of a blower THE SLAKING OF LIME the used air which collects in the top of a theatre or the gases given off from a lime kiln, through our mix- ture of 56 pounds of quicklime and 100 pounds of water. Finally there comes a stage when our breath bubbled through the clear solution no longer causes turbidity. Now should we distill the water and weigh it, provided none has been accidentally lost in the operation we find that it weighs 100 pounds. Obviously something has replaced the 18 pounds of water which in the first instance was held in chemical combination by the quicklime. Furthermore, our white residue now weighs 100 pounds and chemical tests show us that this substance is identical with the limestone which we started with. We know that the carbon dioxide in the expired breath has combined with the hydrated lime to form limestone. Then we can now call limestone carbonated lime. These are examples of chemical changes. We have learned that no chemical change takes place without a corresponding physical change and also that these chemical compounds unite in definite proportions. Having once learned by experiment how much of one substance is held in chemical union with another we are able to predict before we put these substances together just what will happen and to calculate how much substance we shall have at the end of our experiment. SUMMARY OF CHAPTER I. A physical change means a change in form, as the grinding of limestone. The freezing of water, and its conversion into steam are examples of physical changes. SUMMARY OF CHAPTER I 23 In a physical change the form or state of a substance is changed while the intrinsic chemical properties are not altered. A chemical change means a change in composition. It was shown that the heating of limestone causes a chemical change: a gas (carbon dioxide) is given off and the substance left is lime which has physical and chemical properties totally different from limestone. When water is added to lime a chemical change takes place; a new compound possessing new properties is formed. Heat is given off. In all chemical changes heat is either absorbed or given off. In any chemical change there is always a physical change; the reverse is never true. Pure chemical substances contain definite elements in definite proportions. CHAPTER IT. WEIGHTS AND MEASURES. IN our experiments with quicklime and water we used such quantities stated in such terms as we can readily appreciate. As a matter of fact in order to obtain absolutely accurate results we must use very small quantities. We used 100 pounds so that we should not be confused by decimals. In chemical work the metric system is used in all countries. One kilogram (abbreviation kilo or k) is equivalent to 2.2 pounds. One thousandth of a kilogram, a gram, is used as the basis of the metric system; about 454 grams equal 1 pound. An ounce avoirdupois equals about 28.5 grams. Fifteen to twenty drops of pure water weigh 1 gram, and the volume occupied by this amount of water is equal to 1 cubic centimeter (1 c.c.). Decigram, one tenth of a gram, is rarely used, but milligram 0.001 gm. is convenient for stating the doses of the more active medicinal agents. Sixty-five milli- grams (0.065 gm.) are equivalent to 1 grain. . CHAPTER III. THE METALS. A CLASS of substances represented by gold, silver, copper, tin, lead, iron, nickel, quicksilver, etc., pos- sessing the power of conducting heat and electricity, capable of being fused, moulded and of being drawn out into various shapes, and having a peculiar luster, are known to us as metals. Some of them (gold, silver, lead, copper, quicksilver, etc.), occur in comparatively large quantities as metals in nature and they also occur in combinations with other substances as ores. With one exception (quicksilver) at ordinary tempera- tures they are solid. Rusting of the Metals. Our most common example of a metal is iron. This metal is found free in nature in very small quantities. Pure iron has a grayish- white color which on exposure to damp air soon changes to black and then to red. In the process of color change, physical properties of the metal change: if this entire piece is rusted throughout it loses its strength and crumbles, it can no longer conduct electricity or be hammered into various shapes. Its weight increases. The rusting process can be hastened by heating in the air. The Air and the Rusting Process. A piece of iron placed in a glass tube, the air exhausted and the tube 26 THE METALS sealed, will retain forever its luster. Even heat will not bring about the tarnishing or rusting process unless air is admitted. Therefore air or some constituent of air is necessary for rusting. Silver will rust quickly, turning black, when heated in air. Let us, therefore, place some silver filings in a hard glass tube and heat it to redness and allow air to pass through slowly and then to a similar tube containing pieces of untarnished iron. So long as the silver continues to rust the iron will remain untar- nished. Conduct the air, which has passed over the heated silver and iron to a bell jar and when it is collected in sufficient quantities place in it a mouse. The mouse dies very quickly. Now strike a match and thrust it into the bell jar the flame goes out. What have we learned? First, that the substance in the air which causes iron to rust is the same as that which brings about the tarnishing (rusting) of silver; second, that this same substance is also necessary for life; third, that it is also necessary for burning. There- fore, we reason that the rusting of metals, the main- tenance of life by respiration and the burning of wood are closely related if not the same processes. Oxygen. To continue our experiment, let us heat some of the rusted silver and collect the gas given off in the process by conducting it into a bell jar. Now repeat the mouse and the match experiments with this new gas and we get very different results. The mouse is enlivened and the match burns more brightly than in ordinary air. These results lead us to believe that we have gotten that substance in the ELEMENTS 27 air essential for life in a pure state. Other experi- ments show us that this is true. Priestly, an English chemist, who afterward came to America and settled at Northumberland, Pa., discovered these facts by working with rusted quicksilver and called the gas thus obtained dephlogisticated air. Lavoisier later called it oxygen, which means acid producing, thinking (erroneously) that this substance always produces acid. Now we term the rusting of the metals, the burning of substances and the various vital processes in the animal body where oxygen is combined with other substances, oxidations. The new chemical compounds thus produced are known as the oxides. Iron rust is chemically known as iron oxide, etc. Elements. From the foregoing we have learned that there are at least two gases in the air which may be easily taken out, namely, oxygen and carbon dioxide. We know that these substances exist there free or in other words air is a mixture of gases. As has been stated, the oxides and other chemical compounds of metals occur in the earth as ores. Here we have examples of both compounds and mixtures; the oxide of iron is a compound consisting of iron and oxygen in chemical union, whereas it may be mixed with clay or other oxides that may be removed mechanically. Then a mixture may be analyzed mechanically but compounds can be separated into their constituents by chemical processes only. In the process of chemical analysis we arrive at a stage where we can no longer separate the various 28 THE METALS substances into simpler materials. Such a substance, which can no longer be resolved into unlike compo- nents, we term an element. The metals already men- tioned are elements. Brass is an alloy or mixture of two elements, copper and zinc. Oxygen is also an element. Symbols. For simplicity, convenience and economy in writing chemical formula, the names of the various elements are abbreviated or the first letter alone or combined with some other distinctive letter is used; for example, O means oxygen and Os osmium; H means hydrogen and He helium; S sulphur and Si silicon. We hardly see how Fe stands for iron or Cu for copper until we remember the Latin names ferrum and cuprum. Then if an oxide of iron is composed of 1 atom of iron and 1 atom of oxygen we would write it thus FeO; if in another oxide of iron there are two atoms of iron and three atoms of oxygen we write it Fe 2 O 3 . SUMMARY OF CHAPTER III. All substances can be analyzed into their simple components which are known as elements. Metals (iron, gold, silver, copper and mercury) are elements. They cannot be further separated into different sub- stances. Oxygen is also an element. It is a gas occur- ring in the air and is necessary to life. The rusting of metals is a chemical union of the element oxygen with the metals (also elements). The burning of wood is also ah oxidation. Therefore oxidation may take place quickly resulting in the production of heat and SUMMARY OF CHAPTER III 29 a flame or it may proceed quietly and slowly and in the presence of water. In either case oxidation is the same and the total amount of heat given off is exactly the same. In writing chemical reactions symbols are made use of to indicate the various elements. Usually the first letter, capitalized, is used as the symbol for the element, O = oxygen, H = hydrogen. Sometimes the first two letters, as Si = silicon (S = sulphur). Cu = copper (Latin, cuprum) . Ag = silver (Latin, argentum) . CHAPTER IV. MOLECULES AND ATOMS. THE smallest particle of a compound which can exist is called a molecule. The size of molecules vary but they are inconceivably small. Since compounds are made up of two or more elements, then molecules may be still further divided into atoms, which accord- ing to Dalton's theory are the indivisible particles of which all substances are composed. If we take FeO as an example, the smallest particle of FeO which can exist is a molecule, and this molecule consists of an Fe atom and an atom, which are indivisible. Generally atoms do not exist alone but only in combination, that is to say an element for example like oxygen consists of molecules of oxygen each of which is composed of two atoms of oxygen. We do not have then simply O in the air but O 2 , for when an atom of O has no Fe or Cu or other element to combine with, it unites with another atom of O making 2 . The size of an atom is so small that it cannot be determined or even imagined. We know that certain substances give off odors for long periods of time and never diminish in weight. In an attempt to convey some idea as to the size of these minute, infinitesimal particles, Lord Kelvin says: "Imagine a rain drop or a globe of glass as large as a pea, to be magnified up CAN MATTER BE DESTROYED 31 to the size of the earth; each constituent being magni- fied in the same proportion. The magnified structure would be coarser grained than a heap of small shot, but probably less coarse grained than a heap of cricket balls." Can Matter be Destroyed? In our experiments with limestone, we certainly changed this chemical com- pound and reduced its weight by heating. It was found that we drove off a gas and left a white friable substance that would unite with water and in the process would still further suffer chemical and physical changes. But we found that after we obtained this new chemical compound, hydrated lime, we could bubble the gases from the kiln through it and obtain not only the same substance (limestone) with which we started but in exactly the same amount. In other words with careful handling no matter was lost. Con- sider another example: iron rusts and increases in weight and loses its metallic properties. By appro- priate chemical processes we are able to drive away the oxygen and recover in exactly the same amount the metallic iron. Well enough, but how about the burning of coal do we not destroy the coal in this burning process which we have learned to regard as an oxidation? Does not the ash weigh less than the coal did ? The ash does weigh less, for this repre- sents only the mineral parts of the coal. If we had taken the trouble to collect every particle of the gases given off in this combustion (burning) we would have found that the weight of the gases plus the weight of the ash exceeds the weight of the 32 MOLECULES AND ATOMS coal and the sum is exactly equal to the weight of the coal plus the weight of the oxygen used in the process. Therefore nothing is destroyed things may be changed physically and chemically and the products wafted to the four winds of the earth but never de- stroyed. The ashes may be washed away and the gases dissipated but eventually they are gathered into nature's laboratory and rebuilt into combustible materials. If the elements were destroyed all our methods of determining the exact amounts of each present in a compound would be useless. The fact that in chemical reactions there is no change in the weight (mass) of each element is known as the law of the conservation of mass. Energy. The woodsman's method of lighting a fire is to rub' two pieces of wood together until the heat generated is sufficient to start the oxidation of the* wood. By means of the friction mechanical energy was transformed into heat and finally when the heat is present in large quantities another form of energy, namely, light, manifests itself. Heat energy may be transformed into mechanical through the agency of the steam engine and then into electrical energy by the turning of a dynamo. Electrical energy may in turn be transformed into light, mechanical energy or heat by means of incandescent filament, motor and coils, respectively. Just as matter can not be destroyed it is not possible to destroy energy. Energy may be transformed into heat and dissipated (i. e., lost to immediate surround- ings), but it always remains energy, and always exists SUMMARY OF CHAPTER IV 33 in the same amount. In the transformation of the energy contained in coal (latent energy) into steam or electricity (potential or kinetic energy) we may not always finish our experiment with the same amount with which we began but this is due to imperfections in our machinery. The energy still exists somewhere either in a latent or kinetic form. This is known as the principle of the conservation of energy. SUMMARY OF CHAPTER IV. A compound is the result of the union of two elements. The smallest unit of a compound is a molecule. A molecule is composed of atoms. Iron oxide, FeO, con- sists of very small bodies (molecules) of FeO, which in turn are made up of Fe (atom) and (atom). A molecule of an element may consist of one, two, three or more atoms, a molecule of a compound consists of at least two atoms. Matter may suffer chemical or physical change but it cannot be destroyed. In chemical reactions there is no change in the weight (mass) of the sub- stances involved: this is the law of the conservation of mass. Energy undergoes change in form (heat, light, elec- trical current), but is never destroyed. Heat may be radiated and lost to immediate surroundings but it will always exist as energy in some form. CHAPTER V. CHEMICAL PROCESSES. IN raising a weight to a given height a certain amount of energy (kinetic) is required. What becomes of it the weight isn't heated or lighted or electri- cally charged in the process. It is simply stored up as latent energy for when that weight falls and strikes it liberates exactly the same amount of energy that was required to lift it. Now, returning to our limestone experiment, we will remember that in driv- ing off the carbon dioxide it was necessary to use heat; and when we poured water on the quicklime heat was liberated. In the first chemical process heat was absorbed (compare with the raising of the weight), in the second heat was liberated (falling of the stone). In every chemical process heat is either absorbed or liberated. And we have just come to understand why the water boils when it is poured on quicklime. It means that there is a chemical union brought about between the quicklime and water with the liberation of heat. When the lime hydrate is subjected to heat the re- verse is true; that is, the water is driven off and heat absorbed in the process. The amount of heat given off or absorbed in any chemical process is always the same for a given mass of substances. Later we shall learn that our bodies are kept warm by the heat which is liberated during chemical processes. THE LAW OF MULTIPLE PROPORTIONS 35 The Law of Constant Proportions. If the amount of heat absorbed or liberated is always the same for a given mass of reacting substances as we have just learned then we presume that these substances react in a constant ratio to one another. As a matter of fact, this observation simply confirms what had already been found out by different methods. In the early development of chemistry various substances were analyzed in a crude way, but the methods were sufficient to detect the constancy with which the elements occurred in a given chemical substance. For instance, in our lime experiment we started with 100 pounds on heating until constant weight was attained we always had 56 pounds, no more, no less, provided we started with 100 pounds of pure water- free lime. It has been found that whenever quicklime is air slaked, that is, allowed to absorb all the carbon dioxide it will take up, that the combining ratio is always 56 parts of quicklime to 44 parts carbon dioxide. Many other examples can be given to prove that if two substances react chemically and form a third, they enter into combination in a constant proportion. The Law of Multiple Proportions. It has been found that two or three elements may combine to form entirely different substances. Methods for merely detecting the presence of the elements (known in chemistry as Qualitative Analysis) would show no difference between these two compounds, while quan- titative analysis (i. e. } determining the amounts of the elements present) would bring out this difference in composition. Two substances which are used exten- 36 CHEMICAL PROCESSES sively in medicine illustrate this law. Calomel, a white insoluble, non-crystalline, non-poisonous com- pound, is composed of one atom of mercury (quick- silver) and one atom of chlorine (a greenish, pungent gas and a constituent of common salt). The chemical formula for mercury-chloride is HgCl. If instead of one atom of chlorine combined with mercury we have two, namely HgCl 2 , an entirely different chemical substance results. The latter substance is bichloride of mercury or corrosive sublimate, a sol- uble, highly poisonous, crystalline compound, though it contains nothing that is not found in calomel. The difference lies in the relative amounts of each. By quantitative analysis we observe that the number of atoms of chlorine in the second compound is exactly twice the number in the first for every atom of mer- cury. There are no compounds of mercury and chlorine known in which the number of atoms of chlorine is one and a half or three-quarters, etc., times the number found in calomel. There are only two compounds of mercury and chlorine known: if no law governed their combination there would be any number of different substances depending upon the relative amounts of each present at the beginning of the reac- tion. It has been found that when two elements com- bine in more than one proportion, the masses of the one which combine with a given mass of the other bear a simple rational relation to one another. This ratio in the mercury compounds is 1 to 2. When we come to study the oxides of iron, for example, we find several compounds possible: FeO, Fe 3 4 , Fe 2 O 3 . These com- SUMMARY OF CHAPTER V 37 pounds may at first glance seem at variance with our law and our O seems to be present in fractional parts, but let us elevate them to a common iron content as Fe 2 , then our compounds become Fe 6 O, Fe 6 O 8 and Fe 6 9 . The combining ratios of the O is 6, 8 and 9. When we come to study the various combinations of carbon and hydrogen, we find a wider application of this law because more compounds are possible: as CH 4 , C 2 H 6 , C 3 H 8 , C 4 H 10 , etc. The Law of Combining Weights. When chemical substances are subjected to quantitative analysis, the relative weights of the elements entering into a com- pound are found. To use the mercury compounds as example, we find that the ratio of the weight of mercury to the weight of the chlorine in calomel is 199.8 to 35.18, while in corrosive sublimate the ratio is 199.8 to 70.36. We observe that the chlorine factor in the second compound is exactly twice that in the first compound. After analysis of large numbers of compounds the following law has been formulated: Substances combine either in the ratio of their combining weights or in simple multiples of these numbers. SUMMARY OF CHAPTER V. In any given chemical reaction, heat is either absorbed or liberated in an unchangeable ratio to the amounts of the substances reacting. The amount of heat absorbed when 1 gram of lime is made from hydrated lime is exactly the same as that given off when water is poured on 1 gram of lime. If two substances react chemically and form a third, 38 CHEMICAL PROCESSES they enter into combination in a constant proportion. When two elements combine in more than one pro- portion, the masses of the one which combine with a given mass of the other bear a simple rational relation to one another. Substances combine either in the ratio of their com- bining weights or in simple multiples of these numbers. CHAPTER VI. ATOMIC WEIGHTS. THE gas used to inflate balloons is hydrogen. This colorless, odorless gas is so light in weight that it causes the balloon to float in air. Hydrogen burns in oxygen to form an oxide this oxide is water, H 2 O. The gas chlorine of which we have spoken already as being able to combine with mercury also combines with hydrogen to form hydrogen chloride, HC1. The relative weights of the combining substances are: chlorine 35.18 to hydrogen 1; that is, 35.18 grams of chlorine unite with 1 gram of hydrogen to form 36.18 parts HC1. The substance iodine, with which we are all familiar also unites with hydrogen to form hydrogen iodide, HI, and in the ratio of 125.89 I to 1H. We remember that 199.8 parts mercury unites with 35.18 parts chlorine. By continuing our analyses beginning with any one element and finding the ratio of the combining weights of others, then taking these and working to still others we are able to obtain the com- bining weights of all elements in terms of one another. Now, how should we express them? The easiest way is to take the lightest element as our basis and let it equal 1. Since hydrogen is the lightest element we build our system on it stating that the combining weight is 1 and then chlorine will be 35.18; iodine 125.89; mercury 199.8, etc. In the hydrogen chloride combination we 40 ATOMIC WEIGHTS find one in which the smallest quantity of hydrogen has entered. This 1 part of hydrogen must be at least one atom it may be more but it cannot be less. By anal- yzing large numbers of compounds and by certain other methods we arrive at the conclusion that one atom of hydrogen combines with one atom of chlorine. Since the combining weights are respectively 1 (for H) and 35.18 (for Cl) we say that these numbers represent the atomic weights of these elements. In chemical manipulations atomic weights play a very important part and if we can simplify the system of atomic weights to any extent we save the chemist that much. On the basis of hydrogen as 1 nearly all the atomic weights of the elements contain decimals. For example, oxygen when H=l, is 15.88. The accepted system is to let = 16.0 to make it a round number, then H becomes 1.008 and arsenic which was 74.9 on the old basis now becomes 75.0; phosphorus, 30.96 becomes 31.0; mercury 199.8 becomes 200.0. This simplifies a great many calculations. It is usual then to state the atomic weight of oxygen as 16, though it makes no difference which system is used so far as the ultimate results are concerned, for all these numbers are relative. If we begin a calculation with either system we should of course use it throughout the immediate problem. Avogadro's Hypothesis. If one liter of hydrogen and one liter of chlorine are allowed to combine chemically there results not one liter of hydrogen chloride, but two liters of HC1. 1L. H + 1L. C1. = 2L. HC1. Now according to Avogadro's hypothesis "In equal volumes MOLECULAR WEIGHT 41 of all gases, at the same temperature and pressure, there is an equal number of molecules." Therefore, in two liters of HC1 there must be twice as many molecules as there are in one liter of H or one liter of Cl (or any other gas), and since each molecule must contain one atom of H and one atom of Cl it follows that each molecule of H and each molecule of Cl consists of two atoms. Then hydrogen as a gas exists not as H but as H 2 and chlorine as C1 2 . This is in accordance with our previous statement that when atoms of elements have nothing else to combine with they sometimes combine with one another (see page 30). Molecular Weight. The sum of the atomic weights of the elements composing a molecule is called the molecular weight. For example, if a molecular of hydrogen is H 2 and the atomic weight of H = l, then the molecular weight is 2; chlorine exists as C1 2 ; bromine as Br 2 ; nitrogen as N 2 , etc., therefore the molecular weights of these substances are twice their atomic weights. Mercury, however, exists as Hg, while in phosphorus there are three atoms in a mole- cule, P 3 . Further, in substances composed of more than one element as HC1, the molecular weight is the sum of the atomic weight of H plus the atomic weight ofCl, or 1.008 +35.45 = 36.458 = molecular weight of HCL Now for the reason why we learn about atomic weights and molecular weights. When chemical sub- stances react with one another, they join atom to atom not gram to gram. Obviously, if hydrogen is the lightest substance then a gram would contain as many 42 ATOMIC WEIGHTS more atoms as its weight is less than chlorine for instance so that to obtain 36.458 grams HC1 we do not mix 18.229 (one-half of 36.458) grams of H with the same quantity of chlorine but 1.008 gram H to 35.450 Cl. In using gaseous substances we can measure by volume but when we come to work with metals and salts we must weigh the materials. SUMMARY OF CHAPTER VI. The ratio of the units of combining power of elements is spoken of as atomic weight. Because hydrogen is the lightest gas known, it is accepted as a standard and its atomic weight placed, therefore, as 1. Other atomic weights are expressed in terms of hydrogen as 1. The atomic weight of O = 15.88 (i. e., 15.88 times at. wt. of H). Recorded in terms of H many atomic weights contain decimals. To simplify calculations the value of H is placed at 1.008, then at. wt. of O=16.0. Other examples are given in the text. In equal volumes of all gases at the same temperature and pressure, there is an equal number of molecules (Avogadro's hypothesis). Molecular weight is the sum of the weights of the atoms composing the molecule, e. g., mol. wt. of HC1 = at. wt. of H (1.008) + at. wt. of Cl (35.45) = 36.458. If the molecular weight of the compound having the empirical formula CHO 2 is found by physical measure- ments to be approximately 90 and the sum of the atomic weights is 45 (C = 12, H = 1 .008, O = 16) we learn that the real formula of this compound is (CHO 2 )2 or (COOH) 2 . CHAPTER VII. OXYGEN. Occurrence. It will be remembered that our experi- ments already cited in Chapter II proved that oxygen is one of the chief constituents of the air. It was also stated that the various metals occur in nature as oxides; for example, ordinary clay is the oxide of aluminum and silicon. Oxygen also exists in plant and animal bodies and 85 per cent, of the oceans of water is oxy- gen. Atmospheric air contains 23 per cent, pure oxygen. So abundant is the distribution of this ele- ment that it is estimated that one-half of the earth's crust is oxygen. Preparation in the Laboratory. Oxygen in air is mixed with nitrogen, a very inert (i. e., will not combine easily) gas, and it is therefore difficult to separate them. One can decompose water by passing an electric current through it and hydrogen is given off at the positive electrode (where the current enters the water, so-called anode), and oxygen is given off as a gas at the pole where the current leaves, negative or cathode, and can be collected in an inverted tube. The heating of an easily decomposable oxide as mercury oxide is a source of oxygen: 2 HgO = 2Hg + O 2 . (See footnote.) 1 The Latin for mercury is hydrargyrum and its abbreviation, Hg. One will wonder why this equation was not written thus: HgO = Hg + O. In the paragraph on Avogadro's hypothesis and mole- 44 OXYGEN Commercially oxygen is produced by heating equal parts of potassium chlorate and manganese dioxide. The gas given off is purified by washing in different solutions and compressed into tanks. Uses. Tubes or tanks of oxygen are found in every hospital for emergency use in cases of asphyxiation and where there is small working lung space as in pneu- monia or diminished breathing capacity. Oxygen is necessary in nitrous oxide anesthesia. In the arts oxygen is used with illuminating or acetylene gas to produce a flame of intense heating power. With such a flame one can cut through steel plate with the greatest ease. Properties of Oxygen. Oxygen is a transparent gas, possessing neither color nor odor, and weighs one and one-tenth times as much as air. By diminishing the temperature to an extreme degree of coldness (below 119 C.) and exerting very high pressure, oxygen may be liquefied, though with such difficulty that for a long time it was considered impossible. The most familiar chemical phenomenon is com- bustion which has already been referred to as oxida- tion. It makes no difference whether the reactions take place quickly as the burning of wood, or proceeds slowly, as the tarnishing of silver or the rusting of iron, the chemical process is the same, namely, the chemical cules it was stated that certain atoms combine with one another when they have nothing else to combine with. This is true of O, that is, O exists as C>2 though mercury exists as Hg. Then, so far as Hg is concerned, this last equation is possible, but if O must combine with something, it would return to the Hg so soon as it is released unless there is another O present to form Oz. We therefore presume that these processes go in pairs, that is, 2HgO = 2Hg + O. OXIDES 45 union with oxygen. In pure oxygen these reactions take place more easily than in air which is oxygen diluted with an inert gas (nitrogen). For example, a piece of phosphorus exposed to the air will take fire spontaneously and burn quietly, but if placed in pure oxygen, oxidation takes place immediately and with explosive violence. The reason for this is found in the fact that chemical reactions proceed more rapidly as we increase the temperature. When the gas is pure oxygen, all the heat of the chemical reaction goes to raise the temperature of the reacting substances, whereas in a mixture like air a greater part of the heat is absorbed by a non-reacting substance (nitrogen). Oxides. When an element burns or is slowly acted upon by oxygen an oxide is formed. Certain elements are capable of combining with oxygen in different proportions forming totally different compounds. By burning carbon in various amounts of air, we are able to produce a compound containing 1 part of oxygen for every carbon atom and also to prepare an oxide containing two parts of oxygen for every carbon atom. If there is an excess of carbon and very little air the former oxide (CO) called carbon monoxide, results, while with abundant air supply carbon dioxide (CO 2 ) is produced. Carbon monoxide is a deadly poisonous gas while carbon dioxide is relatively harmless. Thus the oxides are named first according to the number of atoms of oxygen present as monoxide, dioxide, trioxide, pentoxide, etc. Where the number of atoms of the element in combination with the oxygen may also be different we use "oits" and "I'c" to indicate oxides 46 OXYGEN poorer or richer in oxygen. FeO is called ferrous oxide, while Fe 2 O 3 is ferric oxide an -ous oxide con- tains less oxygen while -ic oxides contain more. We shall find that this same terminology is applied to other compounds like the chlorides, iodides, sulphides, etc. Ozone. The peculiar pungent odor noticeable in the neighborhood of electrical dynamos and x-ray apparatus is due in part to the presence of ozone in the air. It is best produced by passing a silent dis- charge of electricity through pure oxygen. Ozone is a very active oxidizing agent and for this reason is a very efficient deodorizer and disinfectant. It is used to rid municipal water supplies of bacteria by allowing the water to fall from a tower while ozone is bubbled through it. A few hospitals in America use ozone to sterilize their operating room supplies. Metallic silver when placed in an atmosphere of ozone at ordinary tempera- ture will be covered with a layer of brown oxide; while oxygen under similar conditions will not bring about such an oxidation. By heating to 300 C. ozone is changed to oxygen which in turn may be retrans- formed to ozone by an electric current. On account of this and other facts it is believed that ozone is oxygen plus an extra amount of energy and instead of existing as 2 (oxygen gas) is Os. SUMMARY OF CHAPTER VII. Oxygen is a transparent gas possessing neither color nor odor. It is slightly heavier than air and can be liquefied. Oxygen is necessary for life. It is very SUMMARY OF CHAPTER VII 47 widely distributed in nature (about half the crust of the earth is said to be oxygen bound to metals and other elements forming oxides). Combustion is an oxidizing process, that is, the union of oxygen with any substance. Oxidation may proceed rapidly with the product of light and heat (fire) or very slowly in water solution (life processes). Oxygen may be prepared by decomposing water with an electric current or by heating mercury oxide. Ozone is a pungent gas product by electric sparks in oxygen. There are reasons for the belief that ozone is oxygen plus energy and that the molecule of oxygen gas contains two atoms (O-2) while the ozone molecule contains three atoms (O 3 ). Ozone is used to sterilize water supplies and as an oxidizing agent in the labora- tory. Ozone is also an efficient deodorizer. CHAPTER VIII. HYDROGEN, H 2 . (Mol. wt. = 2.016; At. wt. = 1.008.) Occurrence. Very small quantities of hydrogen occur in the atmosphere of the earth but by spectrum analysis it has been discovered to be very widely dis- tributed in the stars, the sun, and is found in nebulous masses. Deposits of free hydrogen have been found in great salt deposits and in oil wells. Eleven per cent, of water is hydrogen, so that in actual amounts it ranks next to oxygen in the earth's composition. Preparation in the Laboratory. Just as oxygen is given off at the positive pole when an electric current is passed through water, hydrogen the only other com- ponent is given off at the negative 1 electrode and may be collected in the closed end of a tube. The practical method of producing hydrogen is to 1 Since positively charged substances attract negatively charged substances and vice versa we assume that oxygen is a negatively charged substance because it is attracted to the positive pole and that hydrogen is positively charged because it is attracted to the negative pole. We shall later learn that the metals (iron, copper, silver, nickel, sodium potassium, etc.) are like hydrogen in being classed sa positive elements, while chlorine, bromine, iodine, sulphur, fluorine can be classed as negative as a general rule and that one of the positive elements will easily combine with a negative element. Some elements as arsenic and phosphorus combine easily with either oxygen or hydrogen but to attempt to explain this would take us into physical chemistry. PROPERTIES OF HYDROGEN 49 allow hydrochloric acid (the acid which is normally present in the stomach) to act on a metal like zinc: Zn + 2HC1 = Zn C1 2 + H 2 . Uses. Hydrogen is compressed into tanks. It burns with an intense non-luminous flame and hence is used in welding and heating processes. With a hydrogen-oxygen blow-pipe cast iron can be welded. Balloons are inflated with hydrogen because it is lighter than air. In the laboratory hydrogen is used as a reducing agent, that is, used to combine with the oxygen held in combination by some other substance and thus de-oxidize. In this case it is, as a rule, generated in contact with the substance to be reduced for it is more active just at the point of liberation when it is said to be nascent (born). 1 In the bacteriological laboratory hydrogen is used to displace the oxygen of the air when it is desirable to cultivate organisms which will not grow in the presence of oxygen. Properties of Hydrogen. Hydrogen is a transparent gas possessing neither color nor odor and is the lightest of all known substances. It can be liquefied with difficulty by diminishing the temperature to 200 under a pressure of three hundred atmospheres and allowing it to expand into a pressure of about fifty atmospheres when the heat absorbed during the process of expansion will still further cool it until it liquefies. If liquid hydrogen is poured into a test-tube liquid air will flow from the outside and finally the tube will be covered with frozen air. 1 The H atoms are liberated and seek to combine with oxygen before combining with another H to form Hz, 4 50 HYDROGEN Hydrogen combines with chlorine, iodine, bromine to form acids, in fact hydrogen is present in, and is a necessary constituent of, all acids. In combination with nitrogen, hydrogen forms ammonia (NH 3 ) and with carbon forms a long series of so-called organic com- pounds. Hydrogen combines readily with sulphur to form the bad-smelling gas hydrogen sulphide H 2 S, and with palladium and sodium, etc., to form hydrides, Pd 2 H, NaH. Hydrogen seems to possess a special affinity for oxygen. If these two gases are mixed in the ratio of two parts of hydrogen to one of oxygen they form a highly explosive mixture which can be easily set off by a flame. This fact should be remembered when working with hydrogen gas. When hydrogen gas is passed over an oxide of a metal such as iron at high temperatures, the hydrogen combines with the oxygen to form water as in the following equation: Fe 3 O 4 + 4H 2 = 4H 2 O + 3Fe. If water vapor is passed over heated iron the reverse is true: 4H 2 O + 3Fe = Fe 3 O 4 + 4H 2 . How then are we going to know what will happen? It has been found that if hydrogen is present in excess metallic iron will result, but if water vapor is present in excess the oxide of iron is formed. This illustrates the effect of quantity (mass) on chemical reactions and the reaction here given is said to be reversible in that it may proceed either way SUMMARY OF CHAPTER VIII 51 according to the quantity of materials present. To indicate the reversibility of chemical reaction the equation may be written thus: 3Fe + 4H 2 O ^ Fe 3 O 4 + 4H 2 O. Most chemical reactions are reversible. SUMMARY OF CHAPTER VIII. Hydrogen is an element. It is a transparent, colorless, odorless gas and the lightest of all known substances. It can be liquefied. Hydrogen combines with chlorine, iodine, bromine, etc., to form acids; it unites with metals to form hydrides, and with carbon and oxygen to form a long list of organic compounds. It unites with nitrogen to form ammonia (NH 3 ) and with oxygen to form water (H 2 O). Just as the addition of oxygen to a substance is called oxidation the addition of hydrogen is called reduction. When hydrogen is added to a substance containing oxygen, the hydrogen and oxygen unite to form water and split off from the molecule of the sub- stance so that reduction may mean the abstraction of oxygen by hydrogen. If no oxygen is present reduction means simply the addition of hydrogen to the molecule. Hydrogen occurs in small quantities in the at- mosphere and free in the earth's crust. It is present in large amounts in combination: water is 11 per cent, hydrogen (by weight). Hydrogen results from the decomposition of water by an electric current. The volume of the hydrogen given off at the negative pole is twice the volume of the 52 HYDROGEN oxygen given off at the positive pole. Hydrogen is said to be positive because it is attracted to the negative pole. Hydrogen is also prepared by the action of HC1 on zinc. Hydrogen is used in balloons (on account of its light- ness) and in chemistry for reduction. It is used to replace oxygen in certain bacteriological methods. CHAPTER IX. WATER, H 2 O. (Mol. wt. = 18.) WHEN pure hydrogen or any compound containing hydrogen is burned in oxygen or air, water results. If a cold object is held above a flame water condenses on it. This is true if the flame consists of burning gas which has been deprived of all water vapor, so it follows that water is one of the end-products of com- bustion. Occurrence. By far the commonest chemical com- pound is water. Widely distributed as it is over and through the earth's crust, the chief constituent of animal and vegetable matter, it is so common and so familiar that we accept it without inquiry as to its composition. Not until we begin the study of chemistry do we wonder what its makeup is, and since chemistry can tell us about such common things this science no longer seems artificial and set apart from every-day things. From now on we shall look upon water, salt, wood, rocks and coal in a new light that is, from the stand-point of their chemical composition. Three-fourths of the earth's surface is covered with water, 70 per cent, of our bodies is water, and stones which seem to be dry contain an astonishingly large amount of it. Without it the earth would be dead. 54 WATER Uses. The great solvent for chemical substances is water. Without it many chemical substances will not react; for example, an explosive mixture of hydrogen and oxygen will not explode if both gases are perfectly dry; hydrogen and chlorine will remain forever un- combined if mixed in an absolutely dry state. All chemical processes of living matter, plant and animal, take place in the presence of water. Digestion, assimi- lation, oxidation and elimination require water and by the aid of it heat regulation of the body is possible. The decompositions of plants and animals by which the elements are liberated to be reformed into living matter require moisture, since these processes are carried out by bacteria and their ferments which cannot work in the absence of water. Water is the great cleanser and heat regulator of the earth. Properties. The physical properties of water are so well known that it is not necessary to review them here. It is wise, however, to discuss certain facts concerning such a common substance in order that other less common substances and their forms may be compared with it. On account of the fact that water is so easily obtained in a pure condition, it is used as a standard of comparison. The scientific standard of distance is the meter (about 39 inches). One-hundredth of a meter is called a centimeter. The volume of a perfect cube which measures exactly 1 cm. on every side is 1 c.c.; in other words, the amount of liquid which a cubic container measuring inside 1 cm. each way is 1 c.c. From this measurement HYDROMETER 55 of distance we obtain standards for weight measure by using pure water. At 15 centigrade 1 c.c. of water weighs exactly 1 gram, which is the standard for weight measure. (See Weights and Measures, p. 24.) Specific Gravity. (Sp. gr.) Specific Gravity means particular weight. It is a well known fact that a gallon of molasses weighs more than a gallon of water, that is, volume for volume molasses is heavier than water. If 1 c.c. of molasses w r eighs 1.5 gram and 1 c.c. pure water weighs 1 gram then obviously volume for volume molasses is one and one-half times as heavy as water, or accepting water as the standard of com- parison with specific gravity as 1.000 then the specific gravity of molasses is 1.500. Molasses is a solution of sugars in water with some caramel and extractions which add to the color and flavor. The more sugars dissolved in a given volume of water the greater the specific gravity. When we come to study the chemistry of the urine, we shall see that the determination of specific gravity is very important for the reason that we are able to ascertain by a very quick and simple process the amount of solids in the urine. The salts and nitrogen-containing bodies dissolved in it make its specific gravity vary from 1 .020 to 1 .030. High specific gravity leads us to suspect the presence of sugar in urine. Hydrometer. The instrument for quick determination of the specific gravity of liquids consists of a glass bulb filled with air and weighted with mercury so that the 1.000 mark on the graduated stem is exactly at the surface when the instrument floats in pure water. 56 WATER It is called the hydrometer (water-measure) . See chapter on Uranalysis. Lighter Liquids. Ether floats on water and is there- fore lighter than water. 1 c.c. weights only 0.717 gram: its specific gravity is 0.717. Pure alcohol has a specific gravity of 0.797, but when water is added its specific gravity is proportionally increased. By determining the specific gravity of a sample of alcohol one can ascertain how much water has been added. Kerosene and gasoline are obtained from the same source, both are lighter than water, but gasoline has the lower specific gravity. These two liquids are sold on the basis of their specific gravity. If the specific gravity of kerosene is low, it indicates that it contains too much gasoline and is therefore dangerous for use in lamps, and if the specific gravity of gasoline is too high it is not fit for motors. We see that the determina- tion of specific gravity has a very practical bearing, since we may employ it in determining the relative purity of certain liquids. Thermometer. The freezing- and the boiling-point of pure water are the constants upon which the standard chemical thermometer is made. As the name (Centi- grade, 100) indicates the difference between the boiling- point and the freezing-point is divided into one hundred parts and each part called a degree. Then water boils at 100 C. and freezes at C. The only other kind of thermometer with which nurses must be familiar is the Fahrenheit, named after its inventor. Water boils at 212 F. and freezes at 32 F. The difference between the two is 180 F., and it follows that, since BOILING-POINT 57 100 Centigrade covers the same range as 180 F., 1 C. = 1.8 F. If we wish to convert Centigrade to Fahrenheit we multiply by 1.8 and add 32 (because in Centigrade we begin at freezing-point), which equals 0.37 C., means 37 C. above freezing-point, whereas if 37 C. equals in range 66.4 F. we must take it to mean 66.4 above the freezing-point of water. Since 32 F. is freezing-point of water, 66.4 above this point is 66.4+32.0 = 98.4. Therefore 37 C. = 98.4 F. 1 To convert Fahrenheit reading to Centigrade we subtract 32 and divide by 1.8. 2 Boiling-point. It has been stated that the boiling- point of pure water is 100 C. This is true at sea level, but if we go up a mountain we find that the boiling- point gradually becomes less and less. On the top of Mont Blanc the boiling-point is sometimes as low as 84. The reason for this lowering is the decrease in atmospheric pressure. The same result may be ac- complished in the laboratory by reducing the atmos- pheric pressure by means of a vacuum pump. Knowing the boiling-point of water at any place we are able to estimate roughly the height above sea level. If, on the other hand, we increase the pressure on the surface of the water the boiling-point is corre- spondingly raised. This is what happens in a steam boiler or steam sterilizer. In the latter instead -of a 1 37 C. or 98.4 F. is the normal temperature of the human body. 2 Since the Centigrade thermometer is being used more and more in clinical work, and since the confusion of the two systems in carrying out orders might at times be dangerous, it is suggested that the pupil nurses convert several readings, one into the other. Room tempera- ture 68 F. (20 C.), normal body temperature and bath tempera- tures are good points to fix in their minds. 58 WATER temperature of 212 F., we reach 225 F.; but the pressure is also increased to about 15 pounds per square inch. On account of being able to increase the tem- perature above the sea-level boiling-point we are enabled to kill many bacteria not killed by boiling. SUMMARY OF CHAPTER IX. Water is composed of two volumes of hydrogen and one volume of oxygen (H 2 O). By weight 11 per cent, of water is hydrogen and 89 per cent, oxygen. When hydrogen or any compound containing hy- drogen is burned in air (or oxygen) water is one of the products of combustion. Water is necessary for life and for many chemical reactions. It is the great chemical solvent. Water is very widely distributed in nature. It can be easily obtained in a pure state, and is therefore used as a standard of comparison for various physical characteristics. The relative weight of a given volume of any substance compared with an equal volume of water is spoken of as the specific gravity of a substance. The specific gravity of water is taken as 1.0. The standard unit of volume is the cubic centimeter. 1 c.c. weighs 1 gram at 15 C. The instrument for measuring the specific gravity of liquids is the hydrometer. The thermometer registers the degree of heat. There are several kinds; especial mention is made of the Fahrenheit and the Centigrade. Water freezes at C. or 32 F. Water boils at 100 C. or 212 F. To convert Centigrade reading to Fahrenheit multiply SUMMARY OF CHAPTER IX 59 by f (or 1.8) and add 32. To convert Fahrenheit readings to Centigrade subtract 32 and either multiply by f (or 0.555) or divide by 1.8. Boiling-points of liquids decrease with a decrease in pressure and vice versa. Water boils below 100 C. (212 F.) on top of mountains. CHAPTER X. HEAT. Heat Absorbed in Evaporation. The thermometer measures the degree of heat not the amount of heat. Obviously it requires a larger amount of heat to raise a gallon of water 1 than it takes to raise the tempera- ture of a quart of water 1. Just as we have liquid measures and standards of weight we accept an easy and simple standard for measuring heat. The standard unit is the Calorie 1 (large Calorie) which is the heat required to raise the temperature of 1000 grams of pure water 1 C. One thousand grams (1 kilogram) of water equals 1000 c.c. or 1 liter and is approximately one quart, so that a Calorie is approximately the amount of heat necessary to raise the temperature of a quart of water 1 C. or 1.8 F. To raise the temperature of one quart of water from room temperature (20 C.) to the boiling-point (100 C.) about 80 Calories is needed. Now in order to convert this amount of water at 100 C. into steam at 100 C. a great amount of heat is necessary, viz., 540 Calories. This heat necessary to change a substance from a liquid to a gaseous state is known as the heat of vaporiza- tion. The heat of vaporization is absorbed by the steam and given out again on condensation. The heat 1 The small Calorie is the heat required to raise the temperature of 1 gram water 1 C. THE FREEZING OF WATER 61 of vaporization interests us on account of the fact that nature makes use of it in the cooling of the human body; sweat is secreted upon the skin by the sweat glands and the large amount of heat absorbed during vaporization lowers the temperature of the body. The definition of the term Calorie should be remem- bered, for use will be made of it later in the discussion of food values. The Freezing of Water. The well known principle: heat expands, cold contracts, holds good for water between certain limits. On being heated water ex- pands slightly and when the same mass occupies more volume the density or specific gravity becomes less. For this reason specific gravity determinations and accurate measurements of water must, be made at an accepted temperature say 15 C., in order to be comparable. 1 The differences in density thus produced cause up and down currents in bodies of water. 2 Water becomes denser on cooling until the tem- perature of 4 C. is reached and between 4 and it expands, so that when it freezes the ice produced is lighter than water and will float. If water continued 1 Sometimes determinations of specific gravity are made at 25 C. based on the density of water at 25 C., then the result of such a determination of specific gravity of, let us say, a sample of urine would be expressed thus: Sp. gr. 25/25 = 1.0236. This means that the sample of urine at 25 C. is 1.0236 times as heavy as water at 25 C. 2 In the spring and autumn one often notices a peculiar fishy, violet, or aromatic odor in tap water. This is due to the stirring up of algse by the spring or autumn turn-over. These turn-overs result from the difference in density of the top and bottom layers of bodies of water the top cooling, thus becoming more dense settles while the bottom layer comes up. 62 HEAT to contract to ice would form and sink to the bottom. Streams and lakes would freeze throughout. Heat Absorbed in Melting. Ice at cannot change to water at without absorbing a relatively large amount of heat. The amount of heat absorbed by one kilogram of ice (2.2 Ibs.) in changing to water is about 80 Calories, or, the amount of heat abstracted from a liter (about 1 quart) of water in being cooled from 80 C. to zero. Ordinary room temperature is from 20 to 25 C., and refrigerator temperature varies from 4 C, to 14 C., a difference of, let us say, 10 C. Then one kilogram of ice in melting would absorb enough heat to cool over 8 quarts of water from room temperature to refrigerator temperature. Freezing Mixture. When ice is placed in water ice melts and cools the water until there is an equilibrium established between the solid and liquid forms of water at C. When more heat is absorbed more ice goes into the liquid form and the temperature is lowered. If one adds salt to ice or the water-ice mixture it has a tendency to make more ice melt, and when ice melts it absorbs heat; therefore the water-ice-salt mixture becomes colder. If an excess of ice and an excess of salt be present a constant temperature of 18 C. (18 C. below freezing point) is maintained. If the temperature is higher than 18 C. more salt goes into solution and melts more ice which absorbs heat until the constant point ( 18) is reached. If the mixture cools lower than 18 salt crystallizes out of solution and allows ice to form until the temperature is raised to a constant point. SUMMARY OF CHAPTER X 63 The reason for these facts cannot be explained without going into physical chemistry but the applica- tion can be readily seen. This is the method used in making ice-cream or for cooling anything quickly. The essential thing in making such a freezing mixture is to always have present an excess of salt and an excess of ice. It should be remembered that 18 C. is intense cold and that such a mixture should not there- fore be used in an ice-bag applied to a patient. SUMMARY OF CHAPTER X. The thermometer registers the degree of heat. The quantity of heat is expressed in Calories. A small Calorie is the amount of heat necessary to raise the temperature of one gram of water 1 C. The large Calorie (C) equals 1000 small Calories. The large Calorie is generally used. The heat necessary to change a substance from a liquid to a gaseous state (water at 100 C. into steam at 100 C.) is known as the heat of vaporization. Heat of vaporization is exactly equal to heat of condensation. One kilogram of ice at C. absorbs 80 Calories in changing to water at C. Heat of freezing is exactly equal to heat of melting. A mixture of ice-salt-water will maintain a tempera- ture of 18 C. so long as ice and salt are present in excess. CHAPTER XL SOLUTIONS AND PURIFICATION OF SUBSTANCES. SOLUTIONS. WATER is the most important solvent. According to Jones' Inorganic Chemistry over three-fourths of the chemical reactions with which we are familiar take place in water solution; certainly the reactions which take place in nature proceed in aqueous solutions. Nearly all substances solids, liquids and gases are soluble to a greater or less degree in water. Many things are regarded in a practical sense as being in- soluble in water, which are so slightly soluble as to be negligible, nevertheless possess a' definite degree of solubility. We are familiar with the solution of solid things in a liquid but, perhaps, we have not thought of the insolubility of liquids. Alcohol and water are mis- cible in all proportions, that is, the one has unlimited solubility for the other; but chloroform and water will not mix in all proportions. Pour some chloroform into water and it sinks to the bottom. Shake them and the chloroform is emulsified but soon collects again at the bottom of the vessels but in a smaller amount. The water if poured off will retain a definite amount of chloroform and has acquired a sweetish taste and an odor characteristic of chloroform. This is the way EFFECT OF TEMPERATURE ON SOLUBILITY 65 chloroform liniment is prepared. By determining the specific gravity of the solution we find that it is greater than 1, and knowing the specific gravity of chloroform we calculate using the exact figures found, the exact amount of chloroform dissolved. Gases are soluble in water according to the pressure of that particular gas on the surface of the liquid. The presence of other gases creating an enormous pressure would have no effect on this. Heat some water from the tap in a glass beaker and observe the bubbles coming off long before the water reaches the boiling point. This is the dissolved air being driven off. Taste some water which has been recently boiled and observe its flatness. Agitate it with air and it regains its normal taste. Effect of Temperature on Solubility. The general effect of raising the temperature of the solvent is to increase its power to dissolve. Therefore if we wish to hasten the preparation of a solution or to make a stronger solution we heat the solvent. As has been stated in the first chapter some salts take up heat when they are dissolved; that is, the solution gets colder as more salt is dissolved and consequently solution is slow. This is true of magnesium sulphate (Epsom salts), so that when making a solution of this chemical it will shorten the process materially if warm water and heat are used. Gases are less soluble as the temperature is raised, therefore solutions of gases are kept cold. In effer- vescing magnesium citrate the carbon dioxide is in solution in the bottle on account of the pressure. 5 66 PURIFICATION OF SUBSTANCES When the bottle is opened and pressure released the gas comes out of solution causing the foam. If the citrate solution is warm practically all the gas is lost, while in cold solutions more gas is retained. PURIFICATION OF SUBSTANCES. Crystallization. The difference in degree of solubility is made use of in the separation and purification of salts. An earth containing mixtures of salts is leached, the solutions filtered and concentrated to the point of crystallization. When a crystal forms its tendency is to come out pure so that by repeated crystallization salts may usually be freed of their impurities. Water is used extensively as the solvent in the purification of substances, but when the salts have about the same degree of solubility in water, other solvents are made use of. Water itself can be purified by successive crystalliza- tion. In the changing of water to its solid form (freez- ing), ice separates out purer than the mother liquor each time. Many substances which either crystallize with difficulty at very low temperatures or not at all are purified by distillation. Distillation. Distillation is the most widely used and easiest method of purifying water 1 for laboratory 1 One distillation is not sufficient to obtain absolutely pure water. The first and last portions are thrown away since they generally contain soluble gases and ammonia. Water distilled in glass con- tains a trace of alkali (lye) and that from block tin stills contains a trace of tin. Certain bacteria grow in distilled water to a slight extent and even if killed by heat they produce ill effects when injected into the body. Therefore for injection purposes (salvarsan, saline, etc.), water should be double distilled and freshly prepared. SUMMARY OF CHAPTER XI 67 purposes, etc. Various oils and alcohol are freed of impurities by this method. In this process the differ- ences in boiling-point are made use of: distillation is allowed to proceed at a given temperature until all the vapor that comes off at that temperature has been recovered and the temperature is then raised to obtain the next fraction. Such a process is termed fractional distillation. Curiously enough, mercury is separated from substances with which it occurs in nature by being distilled, and the same process is used to further purify this metal in the laboratory. SUMMARY OF CHAPTER XL Nearly all substances are soluble to a greater or less extent in water. A large proportion of chemical reactions take place in water solutions. As a rule, heating increases the solvent power of water for liquids and solids and decreases its solvent power for gases. Other things being equal the amount of gas dissolved in a liquid varies as the pressure of that particular gas on the surface. The difference in solubilities of two substances in any given solvent make it possible to separate the substances by recrystallization. Substances may be freed of impurities by recrystallization if the im- purities possess a different degree of solubility. Liquids having different boiling-points may be separated by distillation. Distillation offers another means of purifying substances. CHAPTER XII. NATURAL WATERS CHEMICAL ACTION OF WATER. Sea Water. Owing to its power of dissolving sub- stances, water as it occurs in nature is almost always impure. Sea water contains large amounts of salts which the waters falling on the surface of the earth and seeping through the crust have leached out and brought with them. The air takes up the water in vapor form and leaves the salts. The Dead Sea which has no outlet to the ocean contains a very large quan- tity of salts (about 25 per cent.), so much that the specific gravity is very high and it is impossible for a person to sink in it. Lakes, Rivers, Springs. The amount of impurities in the waters of springs, rivers and lakes depends primarily upon the character of the region through which the water flowed. If a spring is furnished by waters flowing through sandstones and quartz forma- tions it will contain very little impurities, while springs from salt deposits will, of course, contain varying amounts of minerals. Springs coming from very deep sources where the temperature is high, will, when other things are equal, contain more mineral on account of the increase in solvent power due to increase in temperature. Surface waters are liable to contamination by all sorts of impurities chief among them are the excreta of man and animals. Inland HARD AND SOFT WATERS 69 waters containing relatively large amounts of common salt without other minerals in proportion are looked upon by the sanitarian as possibly contaminated by man or animals. Other tests are necessary to confirm this suspicion. Hard and Soft Waters. Waters slowly passing through swamps take up organic acids which increase the solv- ent action. Such waters passing over limestone will take up relatively large quantities of lime and become what are known as hard waters. Hard waters coagulate soap solutions, that is, when soap is used in them they become turbid and if very hard will precipitate the soap in flakes or curds. If the hardness is due to limestone (calcium carbonate) and magnesium carbonate, boiling will drive off the carbonic acid and allow the salts to precipitate and the water is soft. This is called temporary hardness. If the water contains calcium sulphate or iron sulphate (copperas) it is said to be permanently hard. It is plain that the use of hard waters for washing purposes is not economical because so much soap is needed to produce any effect. The remedies generally used to soften waters are: boiling or the addition of borax (sodium biborate). Waters containing salts in any appreciable amounts should not be used for boilers or for steam sterilizers (autoclaves) because of the residue they leave to form a cake. Waters containing organic acids corrode the sterilizers and even hasten the rusting of instru- ments. A small amount of sodium carbonate (washing soda) is added to the water in which instruments are 70 NATURAL WATERS boiled in order to neutralize any acids that may be in the water and thus prevent corroding and rusting. Rain Water. Rain water is soft and one would think it ought to be pure. Rain is the cleanser of the at- mosphere and brings down besides the dust around which rain drops form certain gases like ammonia. Drinking Water (Potable Water). The fact that a water is clear and cold and contains very small amounts of solid matter does not necessarily mean that it should be drunk. One cannot judge the potability of water entirely by its taste, odor, or appearance. Apparently the best water may contain disease-producing germs like typhoid, cholera, etc., so that a water to be safe must conform to certain bacteriological, as well as chemical and physical standards. Fortunately filtra- tion will also remove a large percentage of the bacteria. Mechanical filtration, after the addition of clarifiers, is often used, but the safest filter is slow sand filtration. The latter is a biological as well as a physical and chemical process. Each grain of sand is coated with a gelatinous membrane due to the growth of organisms and the bacteria coming through impinge upon them and are held. An efficient sand -filter removes 97 to 99 per cent, of the germs of the water passing through them. Porcelain filters through which water is forced under pressure are made for home use as well as for the laboratory. Of this type the Berkefeld and the Pasteur are commonly used. They are made in various degrees of fineness and standardized according to the amount of water which will pass through in a given SUMMARY OF CHAPTER XII 71 time under a given pressure. No filter can be said to be safe until it is tested bacteriologically, and then during an epidemic of a water-borne disease even the filtered water should be boiled before being consumed. SUMMARY OF CHAPTER XII. Natural waters are almost always impure. The amount of salts and other impurities in spring, river or lake waters depends upon 'the region through which the water flowed. Sea water and mineral springs con- tain large amounts of salts. Hot springs and deep well waters are liable to contain large amounts of salts. Water containing large amounts of limestone and magnesia is said to be hard. The lime and magnesia form insoluble salts with soaps and no lather is formed until all the salts are precipitated. The use of hard waters for bathing purposes is uneconomical. Waters may be softened by boiling or the addition of borax. Waters containing calcium and magnesium carbonate are said to be in a state of temporary hardness because this property can be removed by boiling. Permanently hard waters contain magnesium, calcium or iron sul- phate. Rain water and distilled water are soft. Waters from swamps passing over limestone are generally hard. Waters containing salts should not be used in autoclaves or boilers on account of their scale-forming properties. The potability of waters cannot be determined solely by appearance and taste. Polluted waters may appear highly desirable for drinking purposes. CHAPTER XIII. COMPOSITION OF WATER. THE composition of a substance is determined by observations from two stand-points analysis and synthesis. By the analysis of water accomplished by passing an electric current through it we find that two elements are obtained. Hydrogen is given off at one pole (negative) in twice the amount in which oxygen is given off at the other (positive) pole. This shows us that at least water is composed of two parts of hydrogen to one of oxygen or H 2 O. But there may be other substances liberated which do not come off as a gas, so we test the water remaining undecomposed in the flask but find no other substances. Then we make a determination of the molecular weight of water vapor by certain physical chemical means and find it to agree with the formula H 2 O. Now, if our work has been correct we ought to be able to synthetize water from these two elements in these proportions. So two parts of pure hydrogen gas and one part of pure oxygen gas are placed in a vessel and exploded by means of an electric spark. Water is formed and all of both gases disappear. We conclude that water is composed of two parts of hydrogen to one part of oxygen. This reaction gives off a great amount of heat and ION I Z AT I ON 73 the compound formed, water, is very stable. It is very difficult to decompose water by heat. In fact water is a very stable compound though it may com- bine chemically with certain substances with apparent ease it reacts as a whole molecule and is not broken up. It combines with oxides to form the hydroxides such as was indicated in the first chapter. Lime which is the oxide of calcium (CaO) is converted into lime hydroxide or calcium hydroxide (Ca(OH) 2 ) by water thus : CaO + H 2 O = Ca(OH) 2 . Water also combines with the dioxide (S0 2 ) or dioxide (S0 3 ) of sulphur to form acids: SO 2 + H 2 O = H 2 SO 3 (Sulphurous acid). and 80s + H 2 O = H 2 SO 4 (Sulphuric acid). Water is probably the most important chemical compound known. It is a by-product in great many reactions in which hydrogen combines with an OH group. For instance, in the union of an acid and a base about which we shall learn later, water as well as a salt are formed: NaOH + HC1 NaCl +H 2 O. Caustic soda. Hydrochloric acid. Common salt. In this case the H of the acid combined with the OH of the base. This OH group is very important in chemistry and should be remembered for use later. It is called the hydroxyl group. lonization. Substances in solution in water do not exist there as closely bound elements but are slightly 74 COMPOSITION OF WATER pulled apart; as: H Cl or Na OH. .The compound is said to be ionized or in loose formation (dissociated). Each component is an ion and termed positive or negative. The extent of dissociation depends upon the dilution, that is, in weak solutions, we find more dissociation or ionization than in strong solutions. When two or moFe salts are in the same solution there is a constant interchange of ions. Suppose, for example, that saltpetre (potassium nitrate KNO 3 ) is in solution with common salt (sodium chloride, NaCl). The former would be dissociated into K NO 3 and the latter into Na Cl. Immediately there would be an interchange with the loose formation of potas- sium chloride, K Cl and of sodium nitrate Na NO 3 . We may imagine a sort of a Virginia reel formation, the girls representing the negative ions and the boys the positive ions. The positive ions would swing his partner then swing some other negative ion, back to his particular negative ion and then to some other and so on so long as all combinations remained in the game. Let us introduce another chemical like lunar caustic (silver nitrate AgNO 3 ). When the positive ion Ag is in loose combination with any N0 3 ion the interchange proceeds as usual but when an Ag ion comes in contact with a Cl ion they form a compound which is insoluble and precipitate from solution. In terms of our illustration this particular combination of Ag and Cl leave the game immediately and this pre- cipitation proceeds until either all the Ag or the Cl is used up. Either may be in excess and use up all the other and precipitation ceases. This fact is made SUMMARY OF CHAPTER XIII 75 use of in the determination of the amounts of either material in solution. If we have a chloride in solution and want to know how much chlorine there is present, we add a soluble salt of silver in excess and the chlorine is precipitated as silver chloride. This is filtered off, washed, and weighed. If we have silver in solution- some soluble chloride is added, and the precipitate treated as above. The question arises: How are we to know which combination will be insoluble and will precipitate? The answer is: We cannot predict what will happen except by experience. We learn accidentally or by trial, for example, that silver chloride is insoluble and that a solution of a soluble silver salt added to a solution of a chloride will form a precipitate and by analysis we find this precipitate to be silver chloride. In working with chemicals one soon learns and easily remembers those things which will not remain in solu- tion together. These elements are known in Materia Medica and prescription writing as incompatibles. Incompatibles. From what we have just learned it is obvious that silver nitrate and sodium chloride cannot be put in the same solution. Any salt of silver is said to be chemically incompatible with a chloride. The nurse, the pharmacist and the physician must remember the incompatible substances with which they have to work. SUMMARY OF CHAPTER XIII. By analysis with an electric current water is found to contain hydrogen and oxygen in the ratio of two 76 COMPOSITION OF WATER volumes H to one volume O. A mixture of 2H+1O can be exploded by a spark. All the gases disappear and water is formed (synthesis). Water therefore must consist of H 2 O or some multiple of H 2 0, as H 4 O 2 , etc. Molecular weight determinations of steam give ap- proximately 18, which we find is the sum of the atomic weights of H 2 and O, and not of H 4 and O 2 . Therefore the formula of water (vapor) is H 2 O. Water is a very stable compound. Water combines with some oxides to form hydroxides (Ca(OH) 2 , NaOH, etc.), and with SO 2 or SO 3 to form acids (H 2 SO 3 = sulphurous acid; or H 2 S0 4 = sulphuric acid). An acid and a base combine to form a salt plus water. The group OH is called hydroxyl. Substances in solution are ionized or slightly dis- sociated. There is a constant and rapid interchange of ions until two ions, which, when united form an insoluble compound. These two ions and similar ions drop out of the field of interaction. Precipitation continues as long as both varieties of ion are present (i. e. y until one kind is used up). Any two substances which cannot exist together in the same solution are said to be incompatible. CHAPTER XIV. HYDROGEN PEROXIDE. WATER is the monoxide of hydrogen H 2 O, while hydrogen peroxide is the dioxide H 2 O 2 . It is water plus one atom of oxygen. Preparation. Hydrogen peroxide does not occur in nature in any appreciable amounts as far as we know. It is perhaps in the atmosphere in very small quantities. It is prepared by treating the peroxide of barium BaO 2 with hydrochloric acid 2HC1. Ba0 2 + 2HC1 = BaCl 2 + H 2 O 2 . It is then purified by fractional distillation (see Chapter X). The ordinary solution on the market is a 3 per cent, solution in water, though a 30 per cent, solution can be purchased. Properties. Pure hydrogen peroxide is a clear color- less liquid very much like water but is one and one-half times heavier (specific gravity = 1.4996). It has no odor, possesses a peculiar, slightly acrid, taste and produces a soapy froth when taken into the mouth. It gives up the extra atom of O on standing, especially in a warm place and exposed to light. It should, therefore, be bottled in amber-colored glass and kept in the ice-box. The solutions of 3 per cent, hydrogen peroxide on the market generally contain a small 78 HYDROGEN PEROXIDE amount (-5- grain to the ounce) of acetanilid to preserve it. Old solutions should not be used. It is safer to replace the stopper with a cotton plug, which allows any liberated gas to escape. Uses. Hydrogen peroxide gives up its oxygen to oxidizable substances in the sense of the following equation: H 2 2 = H 2 + O. It is therefore useful as an oxidizing agent. The 30 per cent, solution is rarely used in medicine; when dropped on the skin it is very caustic. The 3 per cent, solution is non-poisonous and is an antiseptic and for these reasons used in cleaning ulcers and infected wounds. On coming in contact with pus, blood or living tissues it gives up its oxygen which can act in nascent form to kill bacteria and destroy pus. By means of the bouyant effect of the gas it is useful in cleaning remote portions of a wound but should never be used in a deep or punctured wound because the pressure of the gas liberated may drive the infection further up in the tissues. It is useful in diphtheria and other throat infections and as a mouth wash. Hydrogen peroxide is also capable of bringing about the reduction of compounds containing oxygen. The affinity of the atom of O in H 2 O 2 for other atoms of O is so strong that it unites with one atom to form a molecule of oxygen, O2. Catalyzers. Substances like blood, and pus, etc., which decompose hydrogen peroxide are said to con- tain catalases (ferments), which bring about the liberation of the oxygen but do not enter into the SUMMARY OF CHAPTER XIV 79 products of the reaction. A small amount of catalase will decompose an unlimited amount of H 2 O 2 and remain active. Platinum in a finely divided condition, will decompose hydrogen peroxide, and is an example of an inorganic catalyzer. When zinc is acted upon by HC1 with the formation of H, a small amount of stannous (tin), chloride or platinum chloride is put in to accelerate the reaction. It helps the reaction proceed but is not used up a small amount is useful indefinitely. Another example of a catalyzer which is made use of in ordinary life in platinized asbestos as a pocket lighter. Methyl (wood) alcohol in a metal vial, gives off fumes which come into contact with some platinized asbestos. Air is allowed to mix with it and acting under the influence of the finely divided platinum the oxygen begins to combine with the alcohol. This reaction generates heat until there is enough heat to ignite its alcohol fumes into a blaze. SUMMARY OF CHAPTER XIV. Hydrogen peroxide is water plus one atom of oxygen, that is, H 2 2 . It is prepared by treating barium peroxide, BaO 2 with HC1 and distilling. Hydrogen peroxide is an oxidizing agent, easily giving up the extra atom of oxygen in nascent form. It can also act as a reducing agent on compounds con- taining oxygen. It possesses germicidal power. The ordinary solution is 3 per cent. Stronger solutions are 80 HYDROGEN PEROXIDE dangerous. All solutions should be kept in a dark, cool place and preferably in a cotton-plugged bottle. When hydrogen peroxide comes into contact with living tissues it is decomposed by certain ferments called catalyzers. There are also certain inorganic substances like platinum black, which act catalytically on H 2 O 2 . In cleaning wounds with hydrogen peroxide solutions care should be exercised in its use. There is liability of forcing the infecting agents more deeply into the tissues. Hydrogen peroxide should not be employed in dressing deep and punctured wounds. CHAPTER XV. CHLORINE. (At. wt. = 34.45.) Occurrence. Chlorine does not occur in nature free, but in very large amounts in combination with other elements. It is one of the chief elements in sea water, being found there combined with sodium, potassium, magnesium, and others. It occurs also throughout the crust of the earth in soluble and insoluble combinations. As soon as water comes in contact with the soluble salts they are dissolved and carried to the sea. About one-third of the salts of the Dead Sea is sodium chloride (table or common salt), that is, there are seven pounds of sodium chloride in every hundred pounds of the water. Animal bodies contain chlorine, as chlorides in relatively large amounts. Blood contains about 0.8 parts sodium chloride per hundred. It is given off in the urine, sweat and feces in large amounts. Well waters contaminated with sewage even remotely will exhibit evidences of pollution by reason of high chlorine content. Preparation. The laboratory method of producing chlorine is to heat manganese dioxide with hydro- chloric acid. The latter is oxidized to form water and sets free chlorine: MnO 2 + 4HC1 = MnCl 2 + 2H 2 O + C1 2 , 82 CHLORINE An easier method is to allow hydrochloric acid to act on bleaching powder. Commercially chlorine is produced by passing an electric current through sea-water. The sodium chloride is broken up setting free chlorine and hydrogen, and leaving behind soda lye. Uses. Chlorine is a disinfectant and finds use as such in the cleaning of closets, drains, etc., and in the disinfection of excreta from patients having in- fectious diseases. It is a very active element when free, attacking metals and bleaching colors, and for these reasons must be used with care. It is applied in the form of bleaching powder (hypochlorite of lime, which see). In recent years chlorine has been used in the treatment of public water supplies to rid them of infectious organisms. One part of chlorine in 5,000,000 parts of water is sufficient. In this case also it is applied in the form of bleaching powder. Chlorine water is found in the U. S. Pharmacopoeia as Liquor Chlori Compositus, but is very rarely used. Properties. Chlorine is a yellowish-green gas, soluble in water with the production of a yellow solution. It has a characteristic odor and is very irritating to mucous membranes when inhaled. It should be handled with great care and all experiments with it carried out under a hood with good ventilation. It is heavier than air and settles to the bottom of the hood. Chlorine is very active chemically: it combines with almost all elements. Copper, iron, phosphorus, sodium, potassium, etc., burn in chlorine gas just as PREPARATION 83 in oxygen. Most of those mentioned do not have to be ignited; for example, if copper foil is placed in chlorine gas it glows and CuCl 2 is formed. Sodium and potassium unite with moist chlorine almost ex- plosively forming the chlorides. As a general rule chlorides of metals with the exceptions of silver and lead are soluble in water. Hydrochloric Acid. Reference has already been made to the fact that hydrogen gas and chlorine gas combine volume for volume. If equal volumes of these gases in an absolutely dry state be placed together and an electric spark passed through, nothing w T ill happen, but if a minute trace of moisture be present the mixture will explode, forming two volumes of hydrochloric acid gas. The formula for hydrochloric acid is HC1, or some multiple of this as H 2 C1 2 or H 3 C1 3 , etc., because they combine in equal volumes. If now we determine the molecular weight of hydrochloric acid we find it equal to 36.458. Since H = 1.008 and Cl = 35.45, then HiCli must be the formula. If our formula were H 2 C1 2 or more then the molecular weight determi- nations would give higher figures as 72.916 or 109.374, etc. Occurrence of HC1. Hydrochloric acid interests us because it occurs free in the stomach. It is secreted normally by glands in the stomach wall and is necessary in the peptic digestion of meats, eggs, etc. Preparation. Hydrochloric acid is produced on a commercial scale by allowing sulphuric acid to act on a chloride like sodium chloride (common salt). Sul- phuric acid has a stronger affinity for a base than 84 CHLORINE hydrochloric and therefore displaces it in the sense of the following equation: 2NaCl + H 2 SO4 = Na 2 SO 4 + 2HC1. Sodium Sodium chloride. sulphate. We shall later learn how sulphuric acid is produced. Uses. Hydrochloric acid is used extensively in the industries and in chemical analysis and synthesis. In medicine it is administered in dilute form to patients who have an insufficient secretion of acid in the stomach. Properties. As has been stated hydrochloric acid is a gas. We shall come to know it in water solution. Water will absorb the gas until a 39 per cent, solution is formed. This is the concentrated Hydrochloric Acid (sometimes called Muriatic Acid) of commerce. It has a pungent odor and its fumes strongly irritate mucous membranes. It burns the skin and destroys clothing. The dilute hydrochloric acid of the Pharmacopoeia contains 10 per cent, of the gas (1 part of concentrated hydrochloric acid plus 3 parts water). Hydrochloric acid attacks metals such as iron and zinc with the liberation of hydrogen and the formation of the chloride of the metal: Zn + 2HC1 = ZnCl 2 + H 2 . Zinc chloride. All the characteristics of an acid obtain in hydro- chloric acid. It is in order then to discuss briefly the acids in general. Acids. To say that a substance is an acid is to say it is sour (Latin acidits = sour). Besides attacking TEST FOR HYDROCHLORIC ACID 85 metals with the formation of hydrogen as indicated above, acids unite with bases (like soda lye) to form salts and water: NaOH + HC1 = NaCl + H 2 O. Sodium Sodium chloride hydroxide. (common salt). Acids turn blue litmus 1 red. The sour taste of an acid has been found to be due to the ionized hydrogen. In water hydrochloric + acid is dissociated into H Cl. Vinegar is an acid (acetic acid) and its sourness is due likewise to the + dissociated H (acetic acid in water = CH 3 CO 2 H). Acids then are substances possessing a sour taste; chemically we say a compound possesses acid properties which, when dissolved in some dissociating solvent (like water), yields hydrogen ions. Test for Hydrochloric Acid. To detect hydrochloric acid in a solution, we place in it a piece of paper colored with blue litmus (litmus paper), and should it turn red we know that some sort of an acid is present. Having determined that some acid (an ionized H) is present we want to see whether chlorine is the other ion. We remember from the paragraph on dissociation that when a Cl ion comes in contact with a silver ion a precipitate is formed. We therefore add some soluble salt of silver and in solution it is ionized. + Ag -- NO 3 = AgCl + H - NOs. Dissociated Silver silver nitrate, chloride. 1 Litmus is a blue dyestuff obtained by fermenting certain coarsely ground lichens. It is used to test for acids, turning red when they are present and are turned back to blue by bases. When acids are completely and exactly neutralized by bases, the salts formed have no effect on red or blue litmus. 86 CHLORINE The AgCl comes down as a white precipitate which turns dark on standing. It is soluble in ammonia. We added silver nitrate and obtained a precipitate of silver chloride, therefore there must have been an ion- ized Cl present, and hydrochloric acid was in solution. Bleaching Powder. When chlorine is passed into slaked lime (calcium hydroxide Ca(OH) 2 ), calcium hypochlorite and calcium chloride are formed. 2Ca(OH) 2 + 2C1 2 = CaCl 2 + Ca(OCl) 2 + 2H 2 O. Calcium Calcium chloride, hydrochlorite. The mixture is put on the market as bleaching powder, so named on account of its ability to bleach. In slightly acid solutions chlorine is liberated. On account of the very great chemical activity of chlorine colors are bleached and bacteria killed. Reference has already been made (under chlorine) to the dis- infecting power of this substance. Other Compounds of Chlorine. Chlorine combines with oxygen and with sulphur to form a large number of compounds but they are not of interest here. The chlorides of the metals, which are of interest here, will be discussed under the particular metals. SUMMARY OF CHAPTER XV. Chlorine occurs extensively in nature in common salt beds (sodium chloride) and sea water (chlorides of several metals) . It is a very prominent and important mineral constituent of the animal body. Blood con- tains about 0.8 per cent. NaCl. SUMMARY OF CHAPTER XV 87 Chlorine is prepared by treating bleaching powder with hydrochloric acid or oxidizing HC1 with MnO 2 . Commercially it is produced by passing an electric current through brine. Free chlorine is used as a bleaching and as a steriliz- ing agent. One part free chlorine in 5,000,000 parts of water kills disease-producing organisms like the bacillus of typhoid fever. Chlorine is a yellowish-green gas with a characteristic odor. It is very active chemically. Metals like copper, iron, and sodium burn in it, forming chlorides. Hydrogen unites with chlorine to form a very important compound, hydrochloric acid (HC1). Hydrochloric acid is found normally in the human stomach where it is necessary for protein digestion. When there is not sufficient acid in the stomach HC1 is administered by mouth in dilute form. The concen- trated hydrochloric acid of commerce is a 39 per cent, solution of the gas in water. It is prepared by treating a chloride with concentrated sulphuric acid. Hydrochloric acid attacks metals like zinc, forming the chloride of zinc (ZnCl 2 ) and setting free hydrogen gas. (See Preparation of Hydrogen.) The chemical definition of an acid is as follows: an acid is a compound of H with some radical or nega- tive element which when dissolved in water yields hydrogen ions. Likewise, bases yield hydroxyl groups (OH) on dissociation. Acids and bases neutralize one another, producing salts. The test for hydrochloric acid (or a chloride) is the addition of some soluble silver salt. If chlorides are 88 CHLORINE present a white precipitate of AgCl appears. This precipitate can be dissolved by adding a small amount of ammonia. Bleaching powder is formed by passing chlorine gas over water-slaked lime. The mixture of calcium chloride and calcium hypochlorite formed constitute bleaching powder. CHAPTER XVI. BROMINE IODINE FLUORINE. BROMINE (At. wt. = 80). Occurrence. Bromine, like chlorine, never occurs free in nature. It is found combined with metals forming the bromides in salt deposits and in the sea, though not in such quantities as the chlorides are found. Preparation. The bromine salts may be decomposed by an electric current to produce bromine, or sulphuric acid may be used to set free hydrobromic acid (just as hydrochloric acid is produced) thus: 2NaBr + H 2 SO 4 = Na 2 SO 4 + 2HBr. and we may free the bromine of the hydrobromic acid from its H by allowing the H to combine with O through the agency of some oxidizing agents: 2HBr + MnO 2 + H 2 SO 4 = MnSO 4 + 2H 2 O + Br 2 . Uses. Bromine itself finds very little use in medicine. It is used in the laboratory in one of the methods of urine analysis and is useful in synthetic chemistry. The salts, especially sodium, potassium, lithium, and strontium as nerve depressants. Properties. Bromine is a very heavy, dark, brownish- red, mobile liquid, evolving at ordinary temperatures reddish fumes, highly irritating to the eyes and mucous 90 IODINE membranes. Its odor is peculiar and penetrating, resembling chlorine. Its chemical properties are very much like chlorine: in combination with the metals bromides are formed and in union with hydrogen hydrobromic acid results. The silver salt (AgBr) silver bromide is insoluble in water like silver chloride, so that the addition of a soluble silver salt is also a test for bromides. The appearance of the precipitate obtained (AgBr) is indistinguishable by inspection from the precipitate obtained with chlorides. Chemically we are able to detect the difference by adding a few drops of ammonia to the precipitate: if the precipitate is dissolved quickly we know it is silver chloride if it is dissolved with difficulty it is silver bromide. IODINE (At. wt. = 127). Occurrence. Iodine is very widely distributed in nature but occurs in small quantities. It is found in small amounts with the chlorides and bromides in deposits and in the sea. It was first found in sea weed ash and has since been detected in other plants and in the lower and higher forms of animal life found in the sea. It is stated that a certain tropical sponge contains 14 per cent, (of the dry matter) iodine. In mammals it seems to play a very important role in the thyroid 1 gland since it is found there in considerable amount. (In the sheep over 9 per cent, of the dried 1 The thyroid gland is found in the front of the neck. It is an enlargement of this gland that is called goitre. USES 91 gland is iodine.) White blood corpuscles also contain a very small amount of iodine in their composition. Iodine is also found in minute quantities in the air and dust. Preparation. Most of the iodine is extracted as a sodium salt (sodium iodate, NaIO 3 ) from the salt- petre deposits of Chile. Europe and Japan furnish perhaps a third of the world's production (750 tons per year). The iodate is freed of its oxygen by means of a reducing agent, such as sulphurous acid and the iodine set free by means of the sulphuric acid formed : (1) 2NaIO 3 + 6H 2 SO 3 = 2NaI + 6H 2 SO 4 . (2) 2NaI + 2H 2 SO 4 = Na 2 SO 4 + 2HI. (3) 2HI + H 2 SO 4 = H 2 SO 3 + H 2 O + I 2 . These reactions take place in the same operation, and it will be seen that 1 molecule of H 2 SO 3 is reformed, therefore in carrying out such a process only 5 parts H 2 SO 3 are needed instead of 6 parts as indicated in equation 1. Uses. Iodine was discovered a hundred years ago, but was not employed in medicine until 1831. Today it is used more widely in medicine and surgery than any other substance. It is a counter-irritant, parasiti- cide, absorbent, and alterative. It is widely used to disinfect wounds and is finding application more and more for preparing the site of operation and the operator's hands. It is more efficient when the tissues are dry. Iodine is used in chemical synthesis and analysis, and in staining in the bacteriological laboratories. 92 IODINE In combination with hydrogen as hydriodic acid (HI) and as the iodides of sodium, potassium, and lithium, it is employed extensively in treatment. The world's consumption is about 750 tons per year. Properties. Iodine is a heavy blue-black solid, with a metallic luster. When heated these crystalline scales do not melt but vaporize immediately, giving off dark purple fumes. The vapor will deposit in needle crystals on coming into contact with a cool surface; in fact these properties are made use of in the purification of iodine. The process is called sublimation. Iodine is soluble in alcohol and chloroform but very slightly soluble in water. When potassium iodide is present it will go into water solution easily, producing what is known in the laboratory as Lugol's solution. Chemically, iodine is an oxidizing agent like chlorine and bromine; that is, in water solution it unites with the H of the water in the presence of a reducing agent and leaves the O of the water free to combine with this reducing agent. Thus sulphur dioxide (SO 2 ), a reducing agent, is oxidized to sulphur trioxide (SO 3 ) by iodine in the presence of water. Silver iodide is insoluble in water and ammonia. It differs from the chloride and iodide in color; silver iodide is yellow while the other two silver salts are white. Silver iodide is darkened on exposure to light though not as much as silver bromide. This change affected by light is the basis of photography, which will be discussed in the chapter on Silver and its Salts. Iodine possesses the peculiar characteristic of com- bining with starch to produce a blue color which dis- SUMMARY OF CHAPTER XVI 93 appears on boiling and returns on cooling. Iodine will produce the blue color only when it is uncombined ; thus hydriodic acid or potassium iodide, etc., will not blue a solution of starch. This reaction may be used to test for either starch or iodine if iodine is sought in a solution add starch solution if starch be sought add a drop or two of Lugol's solution. FLUORINE (At. wt. = 19). Fluorine is an active element belonging to the same group as chlorine, bromine, and iodine. It never occurs uncombined in nature. It is found as fluor spar (calcium fluoride) which is used as a flux in iron furnaces. Fluorine is prepared in a manner similar to the method used for chlorine, and on account of its activity cannot be kept in glass but is bottled in paraffin. Its chief interest to us lies in the fact that the sodium salt is sometimes used unlawfully as a preservative in foods. Hydrofluoric acid, HF, is used to etch glass. The Halogens. These four elements, chlorine, bromine, iodine, and fluorine form a group known as the halogens (salt-producing). They possess certain chemical and physical similarities of especial interest to the chemist, but have no application here. SUMMARY OF CHAPTER XVI. Bromine is a very heavy, dark, brownish-red, mobile, caustic liquid. It never occurs free in nature. It is found as a salt (e. g., NaBr) in the sea. 94 FLUORINE It is produced by passing an electric current through a solution of a bromide. Bromine is little used in medicine but is employed extensively in synthetic chemistry. In its chemical properties bromine is very similar to chlorine. Iodine is a heavy, blue-black solid with a metallic luster. It does not melt but vaporizes and can be sublimed. It is soluble in alcohol and chloroform. Very slightly soluble in water, though it will dissolve in a solution of KI in water (Lugol's solution). Iodine is found as sodium and potassium iodate in Chile saltpetre deposits and in sea weeds. It is prepared by treating the iodate with sulphurous acid. Iodine is perhaps the most widely used substance employed in medicine and surgery. It is employed as counter-irritant, parasiticide, absorbent, and altera- tive. Its chemical properties are similar to those possessed by chlorine and bromine though it is not so active chemically. Starch solutions turn blue in the presence of free iodine in cold solutions, on account of the formation of starch iodide. Fluorine is the most active of this group of halogens and must be kept in paraffin. On account of its power to attack glass HF is used in etching and forms the basis of diamond inks. The sodium salt is sometimes used illegally as a food preservative. Chlorine, bromine, iodine and fluorine belong to a group known as the halogens (salt-producing). CHAPTER XVIT. SULPHUR. (At. wt. = 32.) Occurrence. Sulphur is found in a free state as brimstone in volcanic areas 1 and in combination in mineral deposits. Fool's gold which one sometimes sees in coal is a combination of sulphur and iron, Fe 2 S 3 (iron sulphide). Sulphides of copper, zinc, lead, etc., also occur in large quantities in the earth's crust. The so-called sulphur springs contain hydrogen sul- phide, H 2 S, in solution. On exposure to oxygen and oxidizing agents hydrogen sulphide deposits sulphur: 2H 2 S + O 2 = 2H 2 O + 2S. Sulphur is purified by distillation. The chemically pure product is obtained by recrystallization from solution in carbon bisulphide. Uses. Sulphur itself is used in various skin diseases due to parasites and other causes. It is sometimes administered internally for its cathartic action, due, doubtless, to the sulphides formed in the intestinal tract, for sulphur itself is inert. Sulphur is used in gun powder, and in the manufacture of certain kinds of matches. 1 Sulphur deposits are found in Louisiana in sufficient quantities to warrant mining. 96 SULPHUR Burning sulphur is used in fumigations where it is desired to kill insects or animals, such as rats. Sulphur dioxide gas results from the burning and this, while not bactericidal, is very efficient for ridding buildings and vessels of insects, etc. Sulphites and sulphuric acid are useful in many industries. Properties. Sulphur is a pale yellow solid, light in weight, melts at a little above boiling-point of water (at 118 C.) and boils at about 450 C. On being melted it becomes darker in color, gradually assuming a dark brown shade. The solid is crystalline or non- crystalline (amorphous). It is possible to obtain easily two different kinds of crystallized sulphur. Not only are the shapes of these two kinds of crystals different, but they are found to contain different amounts of energy. Chemical Properties. Sulphur is relatively inert at ordinary temperatures but on heating will combine with a large number of elements. It combines with hydrogen to form two compounds H 2 S and H 2 S 2 . The former is far more important as it finds extensive use in the laboratory. By the use of hydrogen sulphide the various metals can be separated into groups on account of the differences in solubility of their sulphides. Every laboratory has a hydro- gen sulphide generator. The odor of this gas is extremely disagreeable and in large quantities poison- ous. At least a part of the odor of decomposing flesh is due to hydrogen sulphide. Sulphur combines with oxygen to form a number of compounds, only two of which need be mentioned. VALENCE 97 SO2, sulphur dioxide, combines with water to form sulphurous acid: so 2 + H 2 o = H 2 so 3 , which acid combines with bases to form sulphites. The gas SO 2 is liquefied and put up in cylinders for disinfection purposes. SO 2 can be oxidized to sulphur trioxide S0 3 , a liquid, which in the presence of water becomes H 2 SO 4 , sulphuric acid. H 2 SO 4 is used in the preparation of many compounds and as a dehydrating agent. It is used in storage batteries and in steel pickling processes. Valence. The student has doubtless wondered that two or more compounds of the same elements can exist and also that some compounds can hold more hydrogen in combination than others. It was seen that chlorine held one hydrogen in chemical union (HC1) while oxygen held two (H 2 O). Bromine, iodine and fluorine held one only, but sulphur held two (H 2 S). Later it will be seen that phosphorus can hold three or five. Now we learn that sulphur can hold two atoms of oxygen or three atoms, and since 1O is equivalent to 2H we have sulphur capable of holding an equivalent of 2H, of 4H or of 6H. This property of holding an element in combination is called valence. We accept hydrogen as the standard 1 then since chlorine can hold one H, the valence of chlorine is also 1; likewise the valence of O = 2. 1 1 These valences have been found experimentally and we simply remember them after we become accustomed to dealing with chemical substances. The student who wishes to know more of chemistry will find a discussion of the Periodic Law of Mendele'eff (in any thorough treatise on Chemistry) very interesting. We are able to predict what the valence will be in some cases before an element is discovered. 7 98 SULPHUR Some elements instead of having a non- variable valence may be able to hold different hydrogen equiva- lents according to the conditions. Chlorine usually has a valence of 1, but sometimes it may be 3 or 5, while sulphur is 2, 4, 6 or even 8. The question of valence is very important in the study of the carbon compounds, so-called organic chemistry to be discussed presently. SUMMARY OF CHAPTER XVII. Sulphur is a pale yellow solid, melting at 118 C. and vaporizing at 450 C. It may be amorphous (non-crystalline) or crystallize in two forms. It is found free in nature in volcanic deposits, and purified by distillation or crystallization out of carbon bisulphide. Sulphur is not very active chemically. It can be burned to the dioxide, SO 2 , or trioxide S0 3 . These compounds are soluble in water, forming respectively, sulphurous (H 2 SO 3 ) and sulphuric (H^SCX) acids. Sulphur unites with hydrogen and metals to form sulphides like H 2 S; Na 2 S, etc. It is applied in skin diseases and burned to rid ships, houses, etc., of insects and rats. Burning sulphur is not advisable for fumiga- tion after contagious diseases, as SO 2 has little germi- cidal action. The power to hold atoms in chemical combination is spoken of as valence. The valence of H = 1; of Cl, Na, K, Br, I, etc. = 1, while the valence of O, S, Mg, Pb, etc. =2. An element may possess different SUMMARY OF CHAPTER XVII 99 valences as S, the valence of sulphur may be 2, 4, 6, etc., of Hg may be 1 or 2 (HgCl or HgCl 2 ). There is no rule about valence except the place of the element in the periodic system and this is complicated. We must simply remember the valences of the elements commonly dealt with. CHAPTER XVIII. SODIUM, Na. (At. wt. = 23.) IN daily life and in the hospital we come Into contact more often with compounds of this element than any other. Sodium is the metal of which common salt is the chloride and it is the basic element in soda lye and washing and in cooking soda. Occurrence. Sodium, on account of its chemical activity is never found free in nature. In combination with other elements it is found everywhere. The sea contains large amounts of the chloride; great mines of the chloride are found in Germany and as nitrate it forms the well known Chili saltpetre beds. Dust contains it in detectable amounts and even the dust- free atmosphere of the midocean holds it in small quantities. Sodium chloride occurs in all the tissues of the body and in the blood; there are about eight parts in every thousand. In biblical times sodium chloride was mined from deposits containing other salts. The expression "Salt has lost its savor" refers to lumps from which all the sodium chloride was taken away and some other salt left. It is impossible for pure sodium chloride to lose its savor unless changed chemically. PROPERTIES ,101 Preparation The chloride and ' nitrate , found, in; nature may be purified by cry^taUi^yo^v 'and 'trie' element itself is obtained by passing strong electric currents through molten soda lye (NaOH). Uses. The saline solution of the hospital is a solution of sodium chloride. It is an essential com- ponent of media for growing microorganisms. There are over twenty-five salts of sodium used in medicine. As lye, extensive use is made of sodium in various industries. It is being used more and more in dye and other chemical syntheses. Soap is sodium combined with acids from fats as we shall see when we study fats. Soda is necessary in glass manufacture the silicate of sodium is one of the chief constituents of this useful substance. Properties. For a long time soda lye was thought to be an element, but when Sir Humphry Davy passed an electric current through it a bright metallic sub- stance rose to the top and took fire when it came in contact with air. This was metallic sodium, a soft, steel-gray solid, which rapidly turns dark in air if a trace of moisture is present. Sodium is so active that it reacts with explosive violence if water is poured on it. Sodium hydroxide is formed and hydrogen is given off: 2Na + 2H 2 O = 2NaOH + H 2 . The heat generated by the reaction may be sufficient to ignite the hydrogen. Sodium combines with chlorine in the presence of moisture to form sodium chloride (common salt). 102 SODIUM Jt>is obtained from sea-water by evaporation or mined from the salt 'deposits and purified by recrystallization. It crystallizes in characteristic hopper-shaped cubes which are hollow in the centre. These crystals melt at 780, and fly to pieces (decrepitate) when heated. It is soluble in water to the extent of 36 parts per 100; i. e., 100 parts saturated solution of sodium chloride contains 36 parts of the salt. Heat increases the solubility very little. Sodium chloride is necessary for bodily function, and therefore a necessary constituent of the food. In flesh food it is present in sufficient quantities to balance the potassium, calcium, and magnesium salts, therefore carnivorous animals do not require salt. On the other hand, plants contain more potassium than sodium, and herbivorous animals seek salt licks in order to maintain the -balance. Man subsisting on a mixed diet requires salt. To much salty food without organic acids (vinegar, fruit juices) has the tendency to cause scurvy. Sodium chloride is the starting-point in the manu- facture of sodium salts. If ammonia (NH 3 ) is added to a saturated solution of sodium chloride and carbon dioxide (COz) passed through it at first ammonium carbonate is formed: NH 3 + H 2 O = NH 4 OH. Ammonium hydroxide. NH 4 OH + C0 2 = (NH 4 )HC0 8 . Ammonium acid carbonate. COOKING SODA 103 Now in the presence of NaCl there is a travel of ions, so we have a formation of ammonium chloride (NH 4 C1) and of sodium acid carbonate, and since sodium acid carbonate is insoluble in the already saturated solution of salts it is precipitated: (NH 4 )HCO 3 + NaCl = NH 4 C1 + NaHCOs. The sodium acid carbonate is filtered off and purified by recrystallization. Cooking Soda. The sodium acid carbonate just described above is what we know as cooking soda. When an acid is added to it carbon dioxide gas is given off: NaHCOs + HC1 = NaCl + H 2 O + CO 2 . This is the principle of the raising of dough the gas given off expands and makes the bread light. In practice hydrochloric acid is not used because: (1) the reaction would take place before the bread is heated and all the gas would come out of the dough leaving it flat; (2) the strong acid would affect the other ingredients; and (3) the resulting compound, NaCl, would make the bread too salty. For these reasons a weak acid is used so that the reaction does not take place until the dough becomes thick enough to hold the gas. An organic acid like lactic (the acid found in sour milk) or tartaric (manufactured from grapes) is used. The latter is mixed in powdered form with the sodium acid carbonate and sold as baking powder. On heating in the presence of moisture the tartaric acid combines with the sodium acid carbonate to form 104 SODIUM sodium tartrate (Rochelle salt) and C0 2 is given off. 1 Sodium acid carbonate is called sodium bicarbonate, for when it is heated to a high degree it gives off one part of CO 2 and sodium carbonate Na 2 CO 3 remains: 2NaHCO 3 = Na 2 CO 3 + H 2 O + CO 2 . This carbonate is known as washing soda. It has more basic properties than the bicarbonate. This salt crystallizes easily and depending upon conditions combines with one, seven or ten parts of water; that is, we obtain the compounds Na 2 CO 3 .H 2 0; Na 2 CO 3 .- 7H 2 O and Na 2 CO 3 .10H 2 O. These molecules of water are not simply occluded within the crystal but are chemically a part of it though it can be driven off by heat just as the C0 2 is driven off from the molecule of the bicarbonate. 2 Sodium carbonate can be obtained also by leaching the ashes of sea weeds. Sodium Hydroxide. If sodium carbonate is treated with lime water Ca(OH) 2 , the interchange of ions would produce the following compounds: Na 2 CO 3 . Ca (OH) 2 . Na 2 (OH) 2 . Ca CO 3 . Now, Ca CO 3 is an insoluble compound, so that it is immediately precipitated and becomes CaCO 3 . Na 2 (OH) 2 is soluble and resolves into the normal 1 This same reaction takes place when Seidlitz powder solutions are poured together, but here the substances being in solution the reaction takes place more vigorously. 2 These molecules of water are known as water of crystallization. This same phenomenon is observed in many other compounds. SUMMARY OF CHAPTER XV I H 105 compound Na OH. The whole reaction is expressed by the following equation: Na 2 CO 3 + Ca(OH) 2 = 2NaOH + CaCOs. The mother liquor after filtration is evaporated and the soda lye (caustic soda or sodium hydroxide, NaOH) is left in white, crystalline lumps. From sodium hydroxide the various salts of sodium are easily made. SUMMARY OF CHAPTER XVIII. Sodium is a soft, steel-gray solid which rapidly tar- nishes in air containing even a trace of moisture. A small amount placed on water reacts violently to form sodium hydroxide (caustic soda) and sets free hydrogen. The hydrogen soon catches fire on account of the heat of the reaction. Hydrochloric acid added to sodium hydroxide forms sodium chloride (common table salt), plus water. Sodium is found everywhere. Even dust contains detectable amounts of this element. Sodium chloride and nitrate occur in large salt beds and the former is the principle salt of sea-water. Sodium chloride is essential to the animal organism. Cooking soda is sodium bicarbonate. Washing soda is more strongly alkaline and is the normal carbonate. Baking powder is a mixture of sodium bicarbonate plus some organic acid. When heated these two sub- stances combine to form a salt and liberate CO 2 which "raises" the dough. Sodium salts are characterized by their easy solu- bility in water. A very minute portion of sodium or a sodium salt imparts an intense yellow color to a flame. CHAPTER XIX. ACIDS AND BASES POTASSIUM. WE have learned that an acid is a sour tasting sub- stance which dissociates in solution into H ions and some other ions. An example of this is hydrochloric acid which gives rise to free H ions and free Cl ions thus : + HCl = H Cl. We come to learn now that a base is a substance which in solution is dissociated, giving rise to free hydroxyl (OH) ions, as sodium hydroxide Na OH. When both an acid and a base are in the same solu- tion they neutralize one another + - + - + - H Cl + Na OH = Na Cl + H.OH. As long as the acid and base are present in such amounts that there are as many H ions as there are OH ions they exactly neutralize one another and the solution is neutral. Molecular Solutions. How do we prepare solutions containing equal numbers of H and OH ions? Let us take as a standard 1 gram of H in a liter of water. But the H is bound up with chlorine if we use hydrochloric acid. Therefore we use such an amount of hydrochloric acid as contains exactly 1 gram H. MOLECULAR SOLUTIONS 107 The atomic weight of H = l, and of chlorine 35.45, making the molecular weight of HC1 = 36.458. There- fore, in 36.45 grams HC1 there is contained 1 gram H. Then 36.45 grams HC1 in 1 liter is our standard solution which we shall term a normal solution. To prepare a solution of a base to exactly neutralize our normal acid solution we must have enough OH ions in solution to be equivalent to 1 gram H in a liter. At. wt. of O = 16 and of H = l, then mol. wt. of OH ion = 17, and we must have, therefore, 17 grams OH to be equivalent to 1 gram H. The at. wt. of Na = 23, OH = 17, mol. wt. NaOH = 40. Therefore, in every 40 grams NaOH we have 17 grams OH or sufficient amount to neutralize 1 gram H ions. To make a normal solution of the base NaOH then we dis- solve 40 grams NaOH in 1 liter of water. One liter containing 36.45 grams HC1 will exactly neutralize 1 liter containing 40 grams NaOH. From the above we learn then that a normal solu- tion of an acid is one which contains 1 gram H ions per liter and a normal solution of a base contains 1 gram OH ions per liter. In the two instances cited above the norlaal solution corresponded exactly with a molecular solution (the molecular weight in grams dissolved in a liter). Suppose we attempt to make a normal solution of sulphuric acid. The formula is H 2 SO 4 and it dissociates into + H - + /S0 4 H/ 108 ACIDS AND BASES that is, two H ions for every molecule. Then a normal solution would be only one-half the molecular solution. An example of a base in which the same is true is cal- cium hydroxide 1 (slaked lime) Ca(OH) 2 , which disso- ciates into + /OH Ca< - X OH. A normal solution is indicated as follows: N/l, twice normal 2N and half -normal N/2. Molecular solutions are labelled M/l, etc., and called molar, half-molar, etc. A chemically normal solution must not be confused with the so-called physiological normal saline solution. The latter is a misnomer and instead of being called normal salt solution it should be called M/8 (eighth molar), for this strength of solution (0.9 per cent. NaCl) has approximately the same osmotic pressure as the blood. Indicators. When acids (which see) were described it was stated that litmus turns red in the presence of acids and blue in the presence of bases. Therefore we have in litmus an indicator which when added to a solution tells us whether the solution is of acid reac- tion or basic (alkaline) reaction. Paper impregnated with litmus is used in urine analysis to ascertain whether the urine is acid or basic. The quickness with which the strip of paper turns indicates the relative strength of the acid or base though only to an approximate 1 Let the student give directions for preparation of N/l HNOs and N/l KOH. POTASSIUM 109 degree. We say then a sample of urine is strongly or weakly acid or alkaline (basic). Volumetric Analysis. Suppose we have a sample of vinegar in which we want to determine the amount of acid. If the vinegar is highly colored we may dilute it and add our indicator. Now we add slowly from a graduated tube (burette) a normal solution of a base until the indicator tells us that the vinegar is now no longer acid but neutral and 1 drop of base added brings about the alkaline color. The basic solution being standard (normal) we know how much base is in a liter and any part of a liter. Also, we can calculate how much acid will be neutralized by any part of a liter of the normal solution of base. Therefore, reading from the burette the number of cubic centimeters of base used to neutralize the acid we can calculate how much acid is in the vinegar. This method of deter- mining the strength of a solution by testing it against a normal .solution is called titration. Titration is the foundation of volumetric analysis. (For further details consult Sutton's Volumetric Analysis.) There are other indicators besides litmus, some of which are acid and some basic, useful according to their particular properties. The nurse will employ, besides litmus, probably only phenolphthalein, alizarin, and Topfer's reagent (dimethyl-amino-azo-benzene). POTASSIUM, K (At. wt. = 39). Potassium is very similar to sodium. It reacts in the same manner and what was said of sodium in general 110 POTASSIUM is true of potassium. It is found as potash (KOH) in ordinary wood ashes and the nitrate is found in large deposits (saltpetre beds). It is found distributed in the human body just as sodium is, though in smaller amounts. Potassium salts are used in medicine like sodium salts; practically the only difference is that they have a more depressing effect on the heart than the latter. Lithium is another metal included in this group with sodium and potassium, called the alkali metals. Little use is made of its salts in modern medicine. SUMMARY OF CHAPTER XIX. An acid in solution is dissociated into H ion + a negative ion (e. g., Cl). A base in solution is dissociated into a metal (let M represent any metal) + OH ion (hydroxyl). The molecular weight of HC1 is 36.458 (H = 1.008; Cl = 35.45). In 36.458 grams HC1 there is 1.008 grams H (approximately 1 gram). A normal solution of an acid is a solution of such an amount of the acid as would represent 1 gram of ionized H per liter, in the case of HC1 36.458 grams per liter. Since masses of acids and bases react molecule for molecule instead of gram for grams, the amount of base necessary to neu- tralize the molecular weight in grams of HC1, would be the molecular weight of the base in grams (NaOH = 40). A normal solution of a base is such an amount of base as would represent 17 grams (sufficient to combine with 1 gram H) of hydroxyl in a liter of water. SUMMARY OF CHAPTER XIX 111 Some acids like H 2 SO 4 possess two available H ions for every molecule therefore a normal solution would be half the gram-molecular weight per liter. This is true also of dihydroxy bases like Ca(OH) 2 . The term normal in speaking of solutions should be confined to strictly chemical meaning. Physiological salt solution is eighth molecular, M/8. Certain dyes like litmus are colored differently by acids and alkalies (bases). The change from a basic to acid reaction is accompanied by a sharp change in the color of the solutions of the dye. This color change is made use of in testing acid solutions against bases for their relative strengths. Such testing or titration (titre = to test) is known as volumetric analysis. The properties and reactions of potassium resemble those of sodium. CHAPTER XX. PHOSPHORUS ARSENIC ANTIMONY- BISMUTH. PHOSPHORUS (At. wt. = 31). THE story of phosphorus forms one of the most interesting chapters in chemistry. Its occurrence in the body, its change of form and varied chemical activity bring it into prominence. Occurrence. Phosphorus occurs as the phosphate of calcium Ca 3 (PO 4 ) 2 in large deposits in the earth and is scattered through the soil also in the form of salts. It does not occur in the free state. The mineral matter of bones is largely calcium phos- phate even to such an extent that bone ash is used as a source for the preparation of phosphorus. The tissues of plants and animals also contain phosphorus in various combinations. Preparation. Calcium phosphate, sand (silicon di- oxide) and charcoal (carbon) are heated in an electric furnace to a high temperature and the free phos- phorus is given off. It is condensed and purified by redistillation. The reaction which takes place is an interchange of silicate and phosphate and a reduction of the phos- phoric acid by the carbon: 2Ca 3 (PO4) 2 + IOC + 6SiO 2 = 10CO + 6CaSiO 3 + 4P. Carbon Calcium monoxide. silicate. PHOSPHORUS 113 Uses. Until recently phosphorus was employed in the manufacture of matches. On account of its poisonous effects on workmen its use now is forbidden by law. Phosphorus is used in rat poison and vermin-killer. In rare instances it is sometimes administered as a therapeutic agent. The hypophosphites were formerly administered as tonics though now it is believed that such treatment is worthless. The pentoxide (P2O 5 ) is a strong dehydrating agent and the chlorides PC1 3 and PC1 5 are used in building up certain organic (carbon- hydrogen) compounds. Properties. Cold phosphorus is a yellowish, brittle solid. On being warmed to room temperature it becomes soft and waxy and melts at a little above body temperature (45 C.). The free elements may exist in four different states: as yellow phosphorus (ordinary form) ; red phosphorus; black phosphorus, and white phosphorus. The first two are more important from a chemical stand-point and more is known of them. Yellow phosphorus is so active that it must be kept under water. When oxygen comes in contact with it, some of the oxygen is converted into ozone, and there is also a chemical union of phosphorus and oxygen in which the oxide is formed, heat is liberated and the phosphorus glows. Undoubtedly the property which phosphorus possesses of glowing in the dark is in some way associated with oxidation. When the temperature reaches 50 C., the phosphorus in contact with oxygen takes fire. Yellow phosphorus is very active chemi- cally, forming some of the compounds which will later 8 114 ARSENIC be discussed. If yellow phosphorus is heated in some inert gas like nitrogen to 250 it is changed to a red amorphous powder known as red phosphorus. In this form phosphorus is not so active. It does not combine readily with other elements and does not take fire when heated in the presence of oxygen. It is much less poisonous than yellow phosphorus. Yellow phosphorus combines with hydrogen to form phosphine PH 3 , a gas, and P2H, a solid. It also combines with oxygen to form oxides P 2 O 3 and P2O 5 , and with hydrogen and oxygen to form acids. Phos- phoric acid H 3 PC>4, which combines with bases to form phosphates and hypophosphorous acid H 3 P0 2 , which in like manner forms hypophosphites, are the more important. A group of important substances found in yolk of eggs, in milk and in brain tissues, known as lecithins contain phosphoric acid, combined with acids from fats and an organic base. ARSENIC (At. wt.= 75). Arsenic in some cases acts like phosphorus and in others like sulphur. It occurs in combinations in nature, never free. Its chief source is the iron com- pound Fe 2 As 3 , and it also occurs in a compound similar to Fool's gold (pyrites). Arsenic is a gray, hard, brittle metal. It combines with various elements: hydrogen, oxygen, sulphur, the halogens and the metals. Its hydrogen compound reminds us of phosphine. Arsine AsII 3 is formed when ANTIMONY 115 hydrogen is generated in a solution containing arsenic. This is the basis of the well-known Marsh test for arsenic. The solution under suspicion as containing arsenic is allowed to run into a flask where hydrogen is being produced by the action of sulphuric acid on zinc. The nascent hydrogen reacts with the arsenic to form arsine which comes out the delivery tubes with the excess of hydrogen. The jet from the delivery tube is lighted and a cold porcelain dish held in the small flame. If arsine is present a metallic film of arsenic is deposited on the dish. Fowler's solution is a solution of potassium arsenite : K 2 HAsO 3 formed by boiling arsenic trioxide with potassium acid carbonate. As 2 O 3 + 4KHCO 3 = 2K 2 HAsO 3 . Arsenic has come into prominence lately on account of its part in the composition of salvarsan which is an elaborated arsenic and benzene compound. Arsenic is a common impurity of mineral acids and certain salts. Special methods are necessary to rid them of the last traces of this metal and the standards set by the Pharmacopoeia allow only minute amounts of arsenic as impurity. ANTIMONY, Sb. (At. wt. = 120). Antimony is very much like arsenic: one seldom attempts to remember its properties except that they are almost the same as arsenic. Even the Marsh test with_slight variation is used as a test for antimony. 116 BISMUTH Our interest in this metal lies in the fact that it is one of the chief components of tartar emetic which is potassium and antimony tartrate. Tartar emetic is put in compound syrup of squills. Antimony is useful in making the alloy used for manufacturing type. BISMUTH, Bi. (At. wt.=208). Bismuth is also like arsenic and closely allied in properties to phosphorus and antimony. .It is not so active chemically as the other members of the group. In fact it is found free in nature. Bismuth is a crystal- lized solid but has not as much of the metallic sheen as arsenic and antimony. Various salts exist. Their formulas are easily pre- dicted when we know that bismuth is trivalent, that is, one atom will hold in combination three atoms of a monovalent element like chlorine. Bismuth hydroxide is Bi(OH) 3 and the nitrate is Bi(NO 3 ) 3 . There is a combination of these two salts called the subnitrate Bi(OH) 2 NO3 in which it will be seen that the bismuth is not entirely nitrated but that two of the nitrate groups (NO 3 ) are replaced by hydroxyl groups (OH). On account of these two hydroxyl groups in the molecule the salt will react basic (alkaline) and is called the basic nitrate. Other metals are capable of forming subnitrates (and other basic salts) also. Bismuth salts are used as astringents and in axray work. If the salts are taken into the intestine and an x-ray made it is found that SUMMARY OF CHAPTER XX 117 the rays are obstructed by the bismuth and shadows are cast on the plate. In this manner the movements and the shape of the stomach and intestines are studied. Sinuses may be studied in like manner. SUMMARY OF CHAPTER XX. Phosphorus is a yellowish, brittle solid which becomes waxy on heating. It may exist in four forms: yellow, red, black or white. Yellow phosphorus is the most common. This variety is the most active chemically. It glows in air and may burn spontaneously if not covered with water. Yellow phosphorus is changed into an inert form (red phosphorus) on heating in an inert gas like nitrogen. Phosphorus occurs as calcium phosphate in mineral deposits and in this form also constitutes the chief inorganic part of bone. The element is obtained from this salt by heating it in an electric furnace with a mixture of sand and charcoal. Phosphorus combines with hydrogen to form phos- phine PH 3 and with oxygen and water to form several phosphoric acids. The salts of these acids are called phosphites, hypophosphites, phosphates, etc. An important class of foodstuffs, lipoids (fat-like substances) may contain phosphorus in combination. Lecithin found in eggs, milk, and brain tissue is a phosphorus containing lipoid. Arsenic is very similar to phosphorus in its com- pounds. It is less active chemically and the free element possesses more metallic properties. It forms arsine with hydrogen (similar to phosphine) and also 118 BISMUTH forms acids and salts, arsenites, arsenates, analogous to phosphorus compounds. The Marsh test depends upon the formation of arsine, AsH 3 . Arsenic com- pounds produce acute and chronic stages of poisoning. Some compounds, potassium arsenite, are adminis- tered, and many of the newer products of chemo- therapy contained arsenic combined with organic radicals (salvarsan). Antimony is similar to phosphorus and arsenic. Potassium and antimony tartrate is the principal medicinal preparation. Bismuth belongs to the same chemical group. It is a metal, sometimes found free in nature. The element is trivalent, i. e., holds in com- bination three chlorine atoms or hydroxyl groups (BiCl 3 , Bi(OH)s). The nitrate Bi(NO 3 ) 3 and especially the subnitrate (Bi(OH)(NO 3 ) 2 , also the sub-gallate are used in local applications for their astringent effect. Bismuth salts are relatively impervious to Rontgen rays and, hence, find use in rontgenography. CHAPTER XXI. CALCIUM. (At. wt.=40.) SODIUM, potassium and lithium are called the alkali metals on account of their ability to form bases (alkalies). Calcium, strontium and barium also pos- sess basic qualities but not to the same extent. This group is called the alkaline earths. Calcium will be discussed as the most interesting representative of the group. Barium is used little in medicine, and only one salt of strontium need be known, viz., strontium bromide. Occurrence. Limestone and chalk are impure cal- cium carbonate. Marble is a pure crystallized lime- stone. The calcium phosphate beds have been men- tioned under phosphorus and fluorite (CaF 2 ) under fluorine. Lime. It has already been stated (Chapter I) that lime results from the heating of limestone. The reac- tion is very simple, carbon dioxide is given off. CaCO 3 = CaO + CO 2 . If the gas be kept confined in the chamber with the lime it will recombine on cooling. This is one of the simplest examples of a reversible reaction. 120 CALCIUM When water is poured on lime slaking takes place with the evolution of heat. . CaO + H.OH = Ca(OH) 2 . In solution calcium hydroxide dissociates into Ca, OH,OH and is therefore basic. Bleaching Powder. Reference to bleaching powder has already been made in connection with chlorine. Chlorine gas passed over moist lime forms equal parts of CaCl 2 and calcium hypochlorite Ca(OCl) 2 or one compound: /oci Ca< cl According to Jones it is the latter toward which most- things point. Whatever the exact composition of bleaching powder may be, it is the most convenient means of trans- porting chlorine, for, all the chlorine which goes into the reaction is recovered when the bleaching solution is rendered acid. Even a weak acid like carbonic suf- fices to replace the chlorine forming calcium carbonate. For this reason bleaching powder exposed to the air over long periods loses its characteristics because of the C0 2 in the air. Plaster of Paris. Calcium sulphate crystallizes with two parts of water of crystallization, CaSO 4 .2 H 2 O. If the crystals are heated slightly above the boiling point of water (107) only a part of the water is lost forming the compound SUMMARY OF CHAPTER XXI 121 This is plaster of Paris. When water is added the compound CaSO 4 .2H 2 O is rebuilt and the crystals fuse together and harden (set) into a mass. If all the water is driven off in the preparation of plaster of Paris (i. e., heated too high) the hydration (adding of water) takes place too slowly to be of use. It will be remembered that plaster of Paris is used in the making of splints. The student now understands why it must be worked quickly after the water is added. Calcium Chloride. Calcium chloride has a great attraction for water and is therefore useful to the chemist as a dehydrating agent. Salts absorbing water on exposure to the air are said to be hygroscopic. Acetylene. Acetylene lamps are well known. Acety- lene is a compound of hydrogen and carbon which burns with a brilliant light. Acetylene is manufac- tured by allowing water to come into contact with calcium carbide CaC 2 . CaC 2 + H 2 O = CaO + C 2 H 2 . Acetylene. Calcium carbide is made by heating a mixture of lime and powdered charcoal in an electric furnace. Flame Tests. A small amount of a calcium salt introduced into the flame of a Bunsen burner produces a dark red color. Strontium salts impart a brilliant red and barium salts a green color to the flame. SUMMARY OF CHAPTER XXL Calcium, strontium and barium have certain char- acteristics in common: their carbonates and sulphates 122 CALCIUM are relatively less soluble in water and all three form hydroxides which are more or less alkaline. These elements are bivalent and constitute a group in the periodic system known as the alkaline earths. The hydroxides of the alkaline earths are not so strongly basic in character as the alkali metal hydroxides. Calcium is the most important member of the group. It does not occur free in nature, but its salts espe- cially the carbonate and phosphate are very common. Marble is crystallized calcium carbonate. Lime is CaO. It results from the intense heating of CaCO 3 . Water slaking means the addition of water to CaO, forming Ca(OH) 2 (calcium hydrate). Bleaching powder Ca< OCl Cl is formed by passing chlorine gas over lime. It is the most economical and convenient method of trans- porting chlorine for bleaching and disinfecting pur- poses. Acids liberate chlorine from bleaching powder. Plaster of Paris is calcium sulphate only partially hydrated. When water is added to a paste of plaster of Paris full hy drat ion quickly results and calcium sulphate crystals (CaS0 4 .2H 2 0) form (setting). Salts like calcium chloride which absorb water from the air are said to be hygroscopic. Acetylene is formed by the action of water on cal- cium carbide. Calcium salts impart a dull red, strontium salts a brilliant red, and barium salts a green color to the Bunsen flame. CHAPTER XXII. MAGNESIUM GROUP. MAGNESIUM (At. wt. = 25). MAGNESIUM, zinc, mercury, cadmium and glucinum belong to this group. Magnesium is of greatest impor- tance, especially from the stand-point of therapeutics. It occurs in nature as the carbonate, the chloride, the silicate, and the sulphate. One of the chief con- stituents of asbestos is magnesium silicate. Epsom salt is magnesium sulphate. The element magnesium, which can be obtained by electrolysis of its salts is a white, very light metal. It burns in air with a brilliant, white light, and for this reason is used exten- sively in fire-works. The Flame Test. Just as metallic magnesium burns characteristically in air other metals like sodium and potassium can be identified by their action in a flame. Sodium burns with a bright, yellow light and potassium gives the flame a peach-blossom color. Many metals, for example, iron, will not burn unless powdered and dusted into a flame, but if a piece be heated white hot and placed in an atmosphere of pure oxygen it will burn brilliantly. In testing the various elements it is not essential that they be in the free state. For example, any salt of sodium will color a flame yellow; 124 MAGNESIUM GROUP potassium salts also yield the characteristic peach- blossom color, strontium salts impart brilliant red; calcium a dull red, barium and copper, green, to the flame. If the light from the flame in which any salt is placed is conducted through a prism in such a manner as to resolve it into its primary colors characteristic bands of different colors will appear. This is known as spectrum 1 analysis. By these means the various ele- ments have been detected in the sun and stars. Magnesia. When metallic magnesium is burned magnesium oxide, MgO, is formed. This compound is what is known as magnesia. It is insoluble in water, but when mixed with water to form a cream it forms the well-known milk of magnesia. Magnesia is generally made by heating the carbonate. The CO 2 is driven off leaving MgO. MgCOs = MgO + CO 2 . Problem for student: What is the chemical reaction when milk of magnesia is taken into the stomach where free hydrochloric exists? MgO + 2HC1= ? Why do you write two parts HC1? Epsom Salt. Epsom salt as found in the Epsom springs, is magnesium sulphate. It crystallizes with seven parts water of crystallization, MgSO 4 .7H 2 O. There is a mineral consisting of MgS0 4 .H 2 O. 1 A beam of white light passing through a prism is spread out into the colors of the rainbow violet, indigo blue, green, yellow, orange, and red. All these colors together are called the spectrum, and the instrument for observing this phenomenon and the various bands formed by different elements is called a spectroscope. The elements in a molten condition produce characteristic lines in the spectrum. MERCURY 125 When magnesium sulphate is dissolved in water it absorbs heat and makes the solution cold, thus reducing the solubility. In making strong solutions, therefore, warm water should be used. MERCURY, Hg. (At. wt.=200). Mercury is the only metal which exists in the liquid state at ordinary temperatures. It occurs free in nature and is separated from the substances with which it is found by distillation. On account of the fact that it solidifies at a very low temperature, and also that its volume is changed to a considerable extent when its temperature is varied, it is very useful for making thermometers. Amalgams. When mercury and gold come into contact they combine the gold loses its yellow color and seems to be silvered over. If there is an excess of mercury the gold will dissolve in it. The same is true of silver, magnesium, calcium, and other metals. The combination of any of these metals with mercury is called an amalgam. Sodium and potassium form amalgams with mercury which are solid at ordinary temperatures and offer a very convenient method of applying the alkali metals for chemical reactions. The amalgamating property of mercury is made use of in the recovery of gold and silver in certain mining processes. Salts of Mercury. Mercury is an example of an element which possesses a variable valence; that is to say, mercury salts exist in which the mercury ion has 126 MAGNESIUM GROUP + valence of one (Hg), there are also those salts in which the mercury ion has a valence of two (Hg). This property of variable valence is best illustrated by the chlorides. Mercury chloride (HgCl) is calomel, a white, non-crystalline insoluble, non-poisonous, powder. Here one readily sees the mercury ion has in combination one ion of chlorine and is therefore monovalent. Mercury bichloride (HgCl 2 ) is corrosive sublimate, a crystalline, soluble intensely poisonous solid. The mercury ion in this compound holds in combination two chlorine ions and is therefore biva- lent. In compounds like HgCl, Hg 2 O, Hgl, etc., where the Hg ion is monovalent the mercury is said to be in the mercurous condition. The compounds mentioned then would be called mercurous chloride, mercurous oxide, and mercurous iodide. Such com- pounds as HgCl2, HgO, HgI 2 are called mercuric com- pounds (mercuric chloride, mercuric oxide, mercuric iodide). The striking difference between mercurous and mer- curic iodide is worthy of comment; mercurous iodide is bright yellow, mercuric iodide is scarlet red. Both mercurous and mercuric compounds are employed extensively in medicine and surgery. The student should be thoroughly familiar with the two chlorides, especially since the substitution of one for the other is a very dangerous error. On standing exposed to light, calomel becomes dark, owing to reduction and deposit of metallic mercury. Obviously at this stage it should not be administered. SUMMARY OF CHAPTER XXII 127 Ammoniated mercury which is the principal con- stituent of Ammoniated Mercury Ointment is a mer- curic compound in which one Cl is replaced by NH 2 . Instead of the formula then is Jig 4 are oxidizing agents. The latter in weak solution is an astringent. CHAPTER XXIV. LEAD SILVER PLATINUM. LEAD (PLUMBUM), Pb. (At. wt. = 207). THE heaviness of lead is proverbial though its specific gravity (11.4) is slightly less than mercury (13.9). Lead is bivalent and forms salts like PbC^. One of the most important salts in medicine is the acetate CH 3 COO\ >Pb CHsCOCK the so-called sugar of lead, useful in skin diseases. Lead is slightly soluble in pure water and more soluble in water in which vegetation has fermented. Since even small amounts of lead taken into the body accu- mulate there and finally cause lead poisoning (plum- bism) it is unsafe to use lead pipes for conducting drinking water. SILVER (ARGENTUM), Ag. (At. wt. = 108). Silver is moderately resistant to chemical action but is attacked readily by nitric acid which converts it into silver nitrate (lunar caustic AgN0 3 ). Silver nitrate is reduced when it comes into contact with organic matter and metallic silver is deposited. This 134 PLATINUM is the reason that one's fingers are blackened by hand- ling it. Silver nitrate forms compounds with albumins, e. g., argyrol, which are used in the treatment of infections of mucous membranes. The halogen compounds with silver, especially silver bromide, darken on exposure to light and for this reason are used to manufacture sensitized plates for photography. When light strikes these plates some slight change is produced according to the intensity and duration of the light. When this plate is put in some reducing agent like pyrogallic acid ("developer") metallic silver is deposited where the light has affected the changes. This deposit of metallic silver will of course form shadows of varying degree. The unchanged silver bromide must be dissolved away before the other light comes into contact with the plate. This is done by washing in sodium hyposulphite, Na 2 S 2 O 3 ("fixing" in "hypo"). The plate is washed again in water and dried. This "negative" is used to "print" the image on sensitized paper which is "developed" and "fixed" in the same manner as the plate. PLATINUM, Pt. (At. wt. = 195). Platinum on account of its usefulness and rare occurrence is worth more than gold. It is very highly resistant to chemical action and for this reason is very useful in chemical procedures. It may be heated in the air without being oxidized, hence the platinum needles with which the bacteriologist transfers cultures. On account of its high melting point and freedom from SUMMARY OF CHAPTER XXIV 135 oxidation, contact points for electrical apparatus are made of platinum. Points and knives for thermo- cautery are also made of platinum. The action of platinum in a finely divided state in bringing about chemical reaction (catalyzer) has already been mentioned. Platinum chloride, PtCU, is used also as a catalyzer; for example, in the produc- tion of hydrogen by the action of HC1 or Zn, a small amount of PtCl 4 greatly accelerates the reaction. Colloids. Platinum is insoluble in water but if a strong current is allowed to pass between two plati- num points in water, some of the metal goes into minute suspension in particles so small that they cannot be seen under a microscope. This is colloidal suspension. Only crystalloids (substances which can be crystal- lized) go into true solution. A colloidal suspension lies between a true solution and a fine suspension. Colloids are precipitated by boiling with acids. Other metals, as silver, gold, copper, etc., can be transformed into the colloidal state. These facts are related to give the student some sort of an idea of what is meant by the term colloid for many of the vital reactions of the cell life are now explained in terms of colloids. SUMMARY OF CHAPTER XXIV. Silver and lead are very similar chemically. Both are fairly resistant to chemical action; nitric acid, however, attacks both, forming nitrates. Sugar of lead is lead acetate Pb(OOC.CH 3 ) 2 . Lead is bivalent. 136 PLATINUM Metallic lead is slightly soluble in pure water and more soluble in slightly acidulated water. Lead poison- ing may result from the consumption of water con- ducted in lead pipes. Silver salts as well as lead are astringent. Silver nitrate and silver albuminates are used in the treat- ment of certain infections. Silver salts find extensive use in photography, because of the fact that light darkens silver chloride, iodide, and bromide. Metallic silver is deposited in proportion to the intensity of the light. See text for description of process. Platinum is very resistant to chemical action and for this reason is useful in many chemical processes. In the bacteriological laboratory the loops for trans- ferring cultures are made of platinum because this metal can be heated to redness so many times with- out deterioration. Thermocautery points and contact points in electrical apparatus are made of platinum for the same reason. Platinum salts (Pt. is tetravalent) like PtCl 4 have the property of stimulating many chemical processes without entering into the final product (catalyzer). Opportunity is taken here to introduce the subject of colloids. An electrical current passed between two platinum points under water will project minute par- ticles of platinum into fine (ultramicroscopic) suspen- sion. Colloidal solutions are not true solutions, but stand somewhere between suspension and solution. The particles in colloidal suspension do not settle out on standing and they cannot be seen with the aid of a SUMMARY OF CHAPTER XXIV 137 microscope. Colloids do not pass through animal membranes as solutions of crystalloids do. Colloids are precipitated by acids. Colloidal chemistry is com- ing into prominence in the study of pathological chemistry. CHAPTER XXV. CARBON, C. (At. wt. = 12.) WE are familiar with the element carbon in non-crys- talline forms as charcoal and graphite. One is sur- prised, however, to learn that the sparkling diamond is nothing more than pure crystallized carbon. The soft, friable, black substance seems to have nothing in common with the white, sparkling stone, the hardest substance known. On complete oxidation both sub- stances yield carbon dioxide and -nothing more. The French chemist, Moisson, was able to produce very small diamonds from charcoal, and more recently larger diamonds have been made. Distribution. When vegetable or animal material is heated in a closed vessel charcoal results, thus showing that carbon enters into the composition of these sub- stances in relatively large amounts. We shall learn later that this element is the chief constituent of living matter. In combination with hydrogen, oxygen and nitrogen it is capable of forming an enormous variety of compounds. For example, the simple substance vinegar is composed of carbon, hydrogen and oxygen, and the highly complex and wonderful animal and vegetable cell substances are chiefly composed of these three elements plus nitrogen. CARBONATES 139 The vast oil and coal deposits are chiefly carbon and hydrogen. Carbon also occurs in minerals (carbonates) . Chemical Properties. Carbon is relatively inert chemically and combines with other elements only under the influence of heat. With lime, for example, under the influence of intense heat, the carbide of calcium Ca 2 C is formed. 1 Heated in the air carbon is oxidized to form carbon dioxide, CO 2 . This is the gas given off by animals during respiration and absorbed by the leaves of plants to be built up into complex vegetable matter. The ultimate product of oxidation of the vegetable and animal matter is CO 2 . When carbon or any of the organic compounds are heated in an atmosphere poor in oxygen more or less of the monoxide (CO) is formed. This gas is deadly to life. In the blood of animals it unites with the hemoglobin (red coloring matter) to form a stable compound and the animal becomes asphyxiated. Many cases of monoxide poisoning have been reported as resulting from the shutting up of stoves over night. Insuffi- cient oxygen is supplied to the glowing carbon and carbon monoxide instead of carbon dioxide is formed. A "flare back" in a Bunsen burner brings about the same condition of insufficient oxygen supply and carbon monoxide is formed. Carbonates. Carbon dioxide in water solution forms the unstable acid H 2 CO 3 , carbonic acid. In the presence of hydroxides of metals the carbonate is formed. For example: 2NaOH + H 2 CO 3 = Na 2 CO 3 + 2H 2 O. Soda Washing lye. soda. 1 When water is added to Ca 2 C acetylene (which see) is formed. ' 140 CARBON When only half the quantity of the hydroxide is present, a hydrogen or acid carbonate is formed. NaOH + H 2 CO 3 = NaHCOs + H 2 O. Soda Cooking lye. soda. The bicarbonate (acid carbonate) of sodium is cook- ing soda to which reference has already been made. The carbonates of magnesium, of calcium, of barium and of strontium are insoluble, therefore the addition of a soluble carbonate to a solution of any salt of the above elements brings about a precipitate. This property is made use of in the separation or estimation of these elements. SUMMARY OF CHAPTER XXV. Charcoal, graphite, and diamond are forms of the element carbon. Carbon is found abundantly in nature : as carbonates it occurs in mineral deposits, and the element enters into a large number of compounds with H, O and N to form the chief constituents of animal and vegetable matter. Carbon is relatively inert chemically. It combines, however, with various elements on heating. Carbon is tetravalent. Completely oxidized it forms, CO 2 , which is given off in expiration. Incomplete oxidation results in the formation of a highly poisonous com- pound, carbon monoxide (CO), which forms a stable union with the red-blood cells. CHAPTER XXVI. COMPOUNDS OF CARBON WITH HYDROGEN. THE study of the various compounds of carbon with hydrogen and with hydrogen and oxygen is known as organic chemistry because the living cells and their products are such compounds of carbon. Marsh Gas. The simplest of the organic compounds is the gas which bubbles up from the stagnant water overlying decomposing vegetable matter. If this gas be collected and mixed with the proper proportion of air it forms an explosive mixture showing that it is an easily oxidizable compound. This is the gas which often collects in mines and explodes when a miner's lighted torch is brought in. By analysis we learn that marsh gas is composed of four parts of hydrogen to one part of carbon. The analyses of other carbon compounds show that carbon combines with the equivalent of four hydrogen atoms. For example, in carbon dioxide, since oxygen has twice the valence of hydrogen, carbon is said to have a valence of four. One atom of carbon will hold four atoms of chlorine in combination, CC1 4 . Since the valence of chlorine is one (remember HC1) thus carbon is again shown to be tetravalent. Many other examples can be cited. 142 COMPOUNDS OF CARBON WITH HYDROGEN Marsh gas is written H H C H H which is our nearest approach to the representation of our conception of the relation of the atoms to each other. In reality we believe that the hydrogen atoms are arranged at the points of a tetrahedron with the carbon atom as the centre. When we increase the number of carbon atoms in a compound one can readily see that the position occupied by the atoms may make a considerable difference in the character of the com- pound. A formula showing the relative positions of the atoms is said to be the structural representation, while the formula indicating simply the relative num- ber of atoms, e. g., CH 4 , is called the empirical formula. H CH 4 H C H Empirical formula. H Structural formula. It will be seen later that two compounds may possess entirely different properties, but have the same empiri- cal formula : only by the structural formula could one distinguish one from another when they are referred to. Methane. Organic chemistry treats of the various compounds built up on the basis of marsh gas as a unit. The chemical name of marsh gas is methane. A mixture of methane and chlorine in diffused day- CHLOROFORM 143 light will react to form chlorine substitution products of methane. CH 4 + C1 2 = HCl + CH 3 C1 mono-chlor-methane. CH 3 C1 + Cla = HCl + CH 2 C1 2 di-chlor-methane. CH 2 C1 2 + C1 2 = HCl + CHCls tri-chlor-methane. CHCla + C1 2 = HCl + CCU tetra-chlor-methane. Chloroform. The third product tri-chlor-methane is the compound familiar to us as chloroform. The structural formula is ci I H C Cl Cl. Chloroform is a heavy, mobile liquid, having a char- acteristic odor and sweet taste. It boils at 62 C., and the vapors are not inflammable as in the case of ether, though when the vapors are brought in contact with a flame they are slightly oxidized, forming carbonyl chloride, a dangerous gas. For this reason care should be exercised in the use of chloroform as an anesthetic near a lamp or gas flame. Commercially, chloroform is produced by the action of bleaching powder on alcohol. lodoform is tri-iodo-methane, that is, iodine is substituted for three hydrogen atoms in methane. CHCls. CNI 3 . CHBr 3 . Chloroform. lodoform. Bromoform. lodoform is a yellow, crystalline solid, with a char- acteristic penetrating odor. Its antiseptic properties are due to the slow liberation of free iodine. 1.44 COMPOUNDS OF CARBON WITH HYDROGEN Methyl Alcohol. If we treat mono-brom-methane with silver oxide in the presence of water we obtain methyl hydrate and silver bromide. 2CH 3 Br + Ag 2 O + H 2 O = 2CH 3 OH + 2AgBr. Methyl hydrate is the compound which we obtain in the destructive distillation of wood and call wood alcohol. Just as the hydroxyl group OH is characteris- tic of a base when joined to a metal, it is when joined to an organic group the characteristic of an alcohol. An alcohol then consists of a hydroxyl group joined to an organic group (radical). Methyl alcohol is represented structurally by the following formula: H I H C OH I H. If R represent any organic radical then an alcohol may be represented by R.OH. Methyl alcohol is lighter than water, and mixes with it in all proportions. This alcohol is used as a solvent in industrial processes and the fumes often cause blindness in the workmen. When taken inter- nally it is a poison and may cause death. On account of its cheapness there is a temptation to use it as an adulterant in the cheaper wines and whiskies which is of course illegal. Formaldehyde. The disinfectant used for fumigat- ing rooms where patients with contagious diseases have been is formaldehyde gas. Formalin is a 40 per cent, solution of the gas in water. An exceedingly small FORMALDEHYDE 145 amount is able to inhibit the growth of bacteria and larger amounts kill them. It is poisonous for animals when taken internally. Strong solutions burn the skin, and there are those who have an idiosyncrasy for it to the extent that even very weak solutions cause violent skin reactions. The chemical constitution of formaldehyde is simple : its structural formula is H H 0=0. We see that this compound is methane in which two H atoms are replaced by O. Observe that the valence of carbon is four and that they are satisfied. In methyl alcohol the O has one valence bound by H, viz.: H H c O H. but in formaldehyde both bonds of the oxygen are attached to the C. Also it will be seen that formalde- hyde has two atoms of hydrogen less. We, therefore, see the relation between methyl alcohol and formalde- hyde. The taking away of hydrogen means oxidation, that is, some oxygen has combined with these two hydrogen atoms to form water which splits off, leaving H The group H c=o, 10 146 COMPOUNDS OF CARBON WITH HYDROGEN (written also CHO) is called the aldehyde group. It is characteristic of the aldehydes because no compound is an aldehyde unless it possesses such a group, and all compounds possessing this group are aldehydes. The aldehydes are made by oxidizing alcohols. In the example here given, methyl alcohol vapor passed over heated copper or platinum wire in the presence of air is oxidized to formaldehyde. If R represent any organic group (radical) then an aldehyde of this radical may be represented by: R.CHO. Organic Acids. If formaldehyde is treated with an oxidizing agent like potassium permanganate KMnO 4 it is oxidized to formic acid : H I H O I I H C=O + O=H C=O. Formic acid is a corrosive, colorless compound, with a penetrating odor occurring in the bodies of ants (Latin formica = an ant) . It also occurs in the hairs of certain caterpillars and in the stings of nettles. It is important to observe that only one hydrogen is replaceable by the metal when the acid is neutralized by a base. The equation for the reaction of sodium hydroxide on formic acid is represented by the following : H Na I J O O I I H C=O + NaOH =H C=O + H 2 O This equation is usually written H.COOH + NaOH = H.COONa + H 2 Q. Sodium formate. SUMMARY OF CHAPTER XXVI 147 Formic acid is the simplest of the organic acids. The carboxyl group (written usually COOH) is characteristic of organic acfds. If R represent any organic radical then an organic acid may be represented by R.COOH. In the case of formic acid R is only H, but in acetic acid (vinegar), for example, R is a methyl group, CH 3 . Now substituting CH 3 for R in our general formula for an organic acid R.COOH, we have CH 3 .COOH (acetic acid). The structural formula of acetic acid may be represented as follows: H I H O I I H C C=O. H SUMMARY OF CHAPTER XXVI. Organic chemistry deals with the compounds of C, H, O, N. The natural compounds made up of these elements are the result of plant or animal life so that the term organic is applied to all such compounds to distinguish them from the mineral or inorganic sub- stances. Marsh gas is the simplest of the organic compounds, and the nucleus about which the more complicated are arranged. It is very important to learn the various 148 COMPOUNDS OF CARBON WITH HYDROGEN substitution products of methane, for this is the foun- dation of organic chemistry. Carbon is tetravalent. The H atoms of marsh gas (CH 4 ) are supposed to be arranged in space about the C atom as a centre. A convenient hypothesis is that each H atom is placed at the angle of a tetrahedron (a four-sided solid). It is conceivable that the figure is equilateral if the four valences of the carbon are satisfied by the same atoms. Such a formula the representation on paper is thus: H I H C H H is called the structural formula, while CH 4 is the empir- ical formula. Two substances may have the same empirical formula but different structural formulas. The four H atoms of methane can be replaced by other monovalent atoms. For example, chloroform is tri-chlor-methane (three H atoms replaced by chlorine). Chloroform is a colorless, mobile liquid. lodoform, a yellow, crystalline solid, is an analogous compound; instead of three chlorine atoms there are three iodine atoms substituted for three hydrogen atoms. If a hydrogen is replaced by a hydroxyl group the result is methyl alcohol, CH 3 OH. An alcohol consists of a hydroxyl group joined to an organic radical, R OH. If two hydrogen atoms are replaced by an oxygen atom the result is formaldehyde, HCHO. SUMMARY OF CHAPTER XXVI 149 The characteristic of the aldehyde group is CHO. Any aldehyde may be represented thus: R CHO. If two hydrogens are replaced by an atom of oxygen, and one hydrogen replaced by a hydroxyl group the result is an organic acid, formic acid, H.COOH. The COOH group is characteristic of an organic acid, R COOH. An acid is an oxidized aldehyde, and aldehyde is an oxidized alcohol, and an alcohol is an oxidized hydro- carbon. H H C H. H Methane. H 1 H C OH. H Methyl alcohol. 3 H ( Formal I >=o. dehyde. OH H C=0. Formic acid. CHAPTER XXVI]. ETHERS. IF methyl (wood) alcohol CH 3 OH is allowed to come in contact with metallic sodium, the alcoholate of sodium is formed: CH 3 OH + Na = CH 3 ONa. We have already seen that sodium has a strong attrac- tion for the halogens (iodine, chlorine, bromine and fluorine), forming with them salts or halides. We have also seen that from methane, CH 4 , a compound of chlorine, bromine or iodine may be formed CH 3 C1, CH 3 Br, CH 3 I. Then if we put either of these three substances in solution with a sodium compound there would be a strong tendency to form NaCl, NaBr or Nal, according to the methyl halide present. If CH 3 .O.Na be brought in contact with CH 3 I, then we would have Nal formed leaving two compounds with unsatisfied or unsaturated bonds : CH 3 .O (or H H C O ) ' . I H and CH 3 (or H C H). H ETHERS 151 The natural result is the joining of these unsaturated bonds thus: H H i I H C O C H. I I H H As a matter of fact this actually happens and we have therefore the equation: CH 3 .O.Na + CH 3 I = CH 3 .O.CH 3 + Nal. Ether. This compound CH 3 .O.CH 3 , written also CH 3 CH is called an ether because of its low boiling-point and elastic property. This (methyl-methyl-ether) is the simplest of the ethers. Later we shall learn that the ether used as an anesthetic (ethyl ether) is of the same type of compound, namely, two organic radicals (R) joined with oxygen: Rs It is not essential that the two radicals are the same the compound R (Ri representing any other radical, for example C 2 H 5 ), is still an ether. 152 ETHERS SUMMARY OF CHAPTER XXVII. Ethers have the constitution R O Ri, in which R and Ri, represent any organic radical as CH 3 , C 2 H 5 , C 3 H 7 , etc. It is not necessary that the two organic groups have the same constitution. The ether given for anesthesia is ethyl ethyl ether, O or C 2 H 6 O C 2 H 6 . Ethyl ethyl ether can be made by the reaction of sodium ethyl alcoholate and ethyl iodide: C 2 H 6 C 2 H 5 .O Na + C 2 H 6 I = Nal + In practice this substance is produced by the action of sulphuric acid on ethyl alcohol. On distillation ether comes over. Ether is a colorless, mobile liquid, having a charac- teristic odor, and a burning, sweet taste. Its specific gravity is about 0.7. It is highly inflammable. It boils at 36 C. Ether for anesthesia contains about 4 per cent, ethyl alcohol and a small amount of water. CHAPTER XXV11I. THE MARSH GAS SERIES. IT has been shown that marsh gas CH 4 in the presence of chlorine, iodine or bromine in the sunlight will gradually form methyl chloride, iodide or bromide. If we heat one of these compounds with sodium, remem- bering the great affinity sodium has for the halogens (Cl, I, Br), we would expect the sodium to combine with the halogen to form a salt. Suppose we take CH 3 I as an example, because this compound is more easily handled than CH 3 C1 or CH 3 Br, on account of the lower boiling-point of the former. If metallic sodium is brought into contact with CH 3 I in ether solution (no water must be present because the sodium would combine with the water to form NaOH) sodium iodide is formed. H H I I H C 1 + Na = Nal + H C . I I H H This leaves an unsaturated compound CH 3 (that is, a carbon atom with only three bonds satisfied). According to chemical laws this compound readily combines with an available body. What, then, is the most available body for this unsaturated compound 154 THE MARSH GAS SERIES to attach itself to? The molecules of substances are so minute that the smallest amount we can appre- ciate must contain millions of molecules. When the reaction takes place between one molecule of each, it suffices to write the equation of single molecules. In the reaction we have millions of molecules of unsatu- rated compounds and as we would expect, they pair off, combining with each other after the manner of the following: 2CH 3 I + 2Na = 2NaI + CH 3 CH 3 . The structural formula of CH 3 CH 3 is H H I H C C H. I I H H This substance is called ethane and is generally written C 2 H 6 . It is a gas similar to methane and found in petroleum and in the neighborhood of oil wells. Chemically it reacts like methane. For example, with iodine each of the H's in turn are replaceable, forming ethyl mono-, di-, tri-, etc., iodide (C 2 H 5 I, C 2 H 4 I 2 , C 2 H 3 I 3 , C 2 H 2 I 4 , C 2 HI 5 , C 2 I 6 ). We come to know C 2 H 5 as an organic radical, ethyl. Alcohol. Then if we treat C 2 H 5 I with silver hydrox- ide we expect to get an alcohol (see methyl (wood) alcohol), according to this equation: C 2 H 6 I + AgOH = Agl + C 2 H 5 OH. This is ethyl alcohol, the ordinary alcohol we know in medicine. SOURCE OF ALCOHOL 155 Properties. Pure alcohol is a colorless, volatile liquid, having an agreeable odor and burning taste. It mixes readily with water and ether in all propor- tions. It is lighter than water, boils at a lower tem- perature (78 C.) and freezes at a much lower tempera- ture ( 130 C.). Alcohol is easily oxidized: fine platinum wire accelerates oxidation of the fumes and will thus set fire to alcohol vapors. It burns in air with a non-luminous, sootless flame. Alcohol is useful on account of its solvent power for oils, resins, and alkaloids. Source of Alcohol. Yeast is a unicellular organism belonging to the fungi. Cultures viewed through a microscope are found to consist of large numbers of single cells. These cells are able to convert certain sugars into carbon dioxide and alcohol. This property is taken advantage of for the commercial production of alcohol. The simplest example is the manufacture of wine in which the natural grape sugar (glucose or dextrose) is changed by the yeast to alcohol and carbon dioxide. The carbon dioxide coming off as a gas gives the appearance of boiling so that the process is called fermentation (fervere = to boil). Wine contains from 10 to 20 per cent, alcohol. In the distillation more alcohol than water goes over on account of the lower boiling-point of the former and the product known as brandy contains about 50 per cent, alcohol. Brandy, of course, contains higher alcohols (fusel oils), which give it the peculiar taste. If lime is added and another distillation carried out the product is almost pure alcohol. Any vegetable containing starch may be used to manufacture alcohol, but the starch must first be 156 THE MARSH GAS SERIES broken down into sugars before the yeast can utilize it. To accomplish this step a ferment is obtained from sprouting barley and this (diastase) added to the cooked starch in the process known as malting. Malt sugar, grape sugar, and dextrins result. Yeast is added and alcoholic fermentation begins. If hops and malt are fermented beer is made. Corn, rye, and potatoes are also used, but here the mash is distilled and whisky is the result. Beer contains from 5 to 8 per cent, alcohol and whisky contains about the same amount as brajidy (40 to 50 per cent.). Fermentations in General. Fermentation is now applied generally to mean the changes brought about by the class of substances known as ferments. The chemical composition of these substances is unknown: they are products of living cells and act according to certain laws. If they enter into combination they are immediately set free, for a small amount of ferment is capable of changing large amounts of substances if given sufficient time. Strictly speaking, fermentation means the destruction of sugars with the production of carbon dioxide and alcohol or organic acids, in contrast to putrefaction, which is the decomposition of proteins with the pro- duction of ammonia and foul-smelling gases. Where microorganisms are capable of inducing fermentation or putrefaction, the former takes precedence over the latter. This means that in a decomposing mixture, as a rule, protein decomposition does not take place and foul odors do not arise until all the sugars are destroyed by fermentation. CHLORAL 157 Acetaldehyde. We remember that formaldehyde HCHO was formed by the oxidation of methyl alcohol, CH 3 OH, or the reduction of formic acid, HCOOH. So acetaldehyde is formed by the oxidation of ethyl alcohol C 2 H 5 OH. H I I II H + H 2 O. Acetaldehyde. Acetaldehyde is written CH 3 CHO. Acetaldehyde is a volatile, colorless liquid with a suffocating odor. It is little used in medicine but its polymer, 1 paraldehyde, is a very useful and safe hypnotic. Paraldehyde. When a drop of sulphuric acid is added to acetaldehyde a condensation occurs. Three molecules of acetaldehyde combine with one another to form one molecule of paraldehyde (CH 3 CHO) 3 , a volatile liquid with a pungent taste capable of pro- ducing sleep with very little depression of the heart or ill after-effects. It should be kept in a cool place. Chloral. This well-known hypnotic is a chlorinated acetaldehyde. On treating acetaldehyde with dry chlorine gas the following reaction takes place: CH 3 CHO + 3C1 2 = CClsCHO + 3HC1. Chloral is then tri-chlor-acetaldehyde. Chloral, itself, is a colorless, oily liquid, but on the addition 1 Polymer is a molecule consisting of two or more molecules con- densed into one. The verb is polymerize. 158 THE MARSH GAS SERIES of water, crystals of chloral hydrate form, CC1 3 .CHO.- H 2 O. This is the compound usually prescribed. Acetic Acid. The next step in the oxidation of ethyl alcohol, after acetaldehyde is formed, is the corresponding acid (compare formic acid). CHsCHO + O = CHsCOOH Acetaldehyde. Acetic acid. and the reverse is true, CHsCOOH + H 2 = CH 3 CHO + H 2 O by the reduction of acetic acid we obtain acetalde- hyde. Acetic acid is the chief constituent of vinegar. The usual method of manufacture is by growing a special fungus or mould in weak alcohol. When apple juice is used fermentation first takes place, and the sugar is changed to alcohol (hard cider) then acidifica- tion begins. Vinegar contains from 1 to 3 per cent, acetic acid. The slimy sediment sometimes seen consists of masses of the mould which forms the vinegar (mother of vinegar). SUMMARY OF CHAPTER XXVIII. The marsh gas series consists of carbon-hydrogen compounds of gradually increasing complexity, begin- ning with methane CH 4 and progressing by the suc- cessive additions of CH 2 . Ethane is CH 4 +CH 2 or C 2 H 6 . Propane = C 2 H 6 + CH 2 = C 3 H 8 , etc. Ethane may be formed from methane by first pro- ducing the mono iodo methane CH 3 I and treating this SUMMARY OF CHAPTER XXVIII 159 compound with metallic sodium in a water-free medium. The marsh gas series may be built up in this manner. The alcohols may be formed by treating the alkyl iodide with AgOH. Ethyl alcohol C 2 H 5 OH is the common (grain) alcohol of commerce. It is a colorless, volatile liquid, lighter than water, miscible in all proportions with water and ether. The commercial source is the fer- mentation of sugars by yeast. Ferments are substances of unknown composition which are capable of bringing about repeated chemical changes without entering into the end-products. A small amount of ferment can change an indefinite amount of sugar if given time enough. The aldehydes of the marsh gas series are made by oxidizing the corresponding alcohol. Acetaldehyde is the result of the partial oxidation of ethyl alcohol. Its formula is CH 3 .CHO. Paraldehyde is formed by the condensation of three molecules of acetaldehyde. Chloral is a chlorinated acetaldehyde; CC1 3 CHO. Acetic acid (vinegar) may be produced by further oxidation of ethyl alcohol or acetaldehyde. It is manufactured by growing a special fungus (mother of vinegar) in weak alcohol. Vinegar contains 1 to 3 per cent, acetic acid. The formula for acetic acid is CH 3 COOH. Let the student compare the formulas of ethane, ethyl alcohol, acetaldehyde, and acetic acid. CHAPTER XXIX. THE PARAFFINS. PETROLEUM is a mixture of a large number of com- pounds composed of carbon and hydrogen. The simplest of these products, methane (marsh gas), we have already described. This has the composition CH 4 , and the next member of the series is ethane which we learned is C 2 H 6 . It will be seen that the difference between these members is CH 2 . Ethane was made from methane by first preparing the iodide CH 3 I and treating this with sodium. By following this method the various members beginning with methane may be prepared: Methane CH4 gas Ethane C 2 H 6 " Propane CaHs " Butane C4Hio liquid Pentane C 6 Hi 2 and so on. The structural formula for butane as an example is: H 3 H. H Observe that each end carbon has three hydrogen atoms, while the included carbons hold only two. THE HIGHER ALCOHOLS 161 In building these compounds a methyl group, CH 3 , is added, but it will be remembered that one of the end hydrogens is split off to make room for the attach- ment of the carbon atom. This explains why the differ- ence between successive members is CH2 rather than CH 3 . As we proceed upward in this series the compounds become less volatile, that is, their boiling-point in- creases. These compounds have been isolated from natural petroleum, but it is a less difficult task to build them from methane and ethane than to attempt to separate them. Gasoline is a mixture of several members of this series. Derivatives of the Hydrocarbons. Since methane and ethane furnished us on oxidation the corresponding alcohols (methyl and ethyl) it is possible to prepare the alcohols of the other members of this series. Hence we have propyl alcohol, butyl alcohol, and so on. We also have the corresponding aldehydes, ethers, and acids. The Higher Alcohols. Ethane we have seen is C 2 H 6 and ethyl alcohol is C2H 5 OH. Then if propane is C 3 H 8 , propyl alcohol is CsHrOH and its structural formula is: H. H This we call a primary alcohol, but if the hydroxyl group, which we remember is the characteristic group 11 162 THE PARAFFINS of an aclohol, is attached to the middle C as in the following : we name it a secondary alcohol. We shall find little need of distinguishing between primary and secondary alcohols, but we shall find it to our advantage to pay close attention to the following compound: OH OH OH I I I H C C C H. I I I H H H Observe that there are three alcohol groups here, and we call it a tri-hydroxy alcohol or tri-atomic alcohol. This is the mon-atomic alcohol already described with the further replacement of two hydrogen atoms by hydroxyl groups. If there were two hydroxyl groups it would be called a di-atomic alcohol. Glycerine. The tri-atomic alcohol, whose formula has just been stated, is the substance known in medicine as glycerine. Properties. Glycerine is so named on account of its sweet taste. As we ordinarily know it, this tri- hydroxy alcohol is an odorless, clear, thick liquid, although in its pure state it is crystalline. It has the property of absorbing water (hygroscopic), and on this account a thin film will keep surfaces moist. It is useful as a vehicle in pharmacy. Glycerine is made from fats by treating them with SUMMARY OF CHAPTER XXIX 163 superheated steam, and it is a by-product in the manufacture of soap. The reactions will be discussed under fats. SUMMARY OF CHAPTER XXIX. The paraffin series consists of saturated hydrocarbons of the marsh gas series. Various members of the series are found in petroleum. The simplest member is methane, CH 4 . Ethane is C 2 H 6 . By the addition of a CH 2 group to any member the next higher is obtained. The first four members are gases the remainder are liquids (excepting the most complex members which are solid at ordinary temperature) . The structural formulas of these compounds are very important. The empirical formulas tell only a very small part of the story of their composition. Derivatives of these hydrocarbons are known. These correspond to the various compounds of methane formed by replacing the H atoms with other elements or groups. It will be seen that more derivatives of the same kind are obviously possible; for example, while in the case of methane only one alcohol was possible, in propane three are possible. Likewise more chlorine derivatives are possible in the higher members of the series. In a primary alcohol the hydroxyl group is attached to the end carbon; in a secondary alcohol this group is attached to the second carbon, etc. A monatomic alcohol possesses one hydroxyl group: a diatomic alcohol, two groups, etc. 164 THE PARAFFINS Glycerine is a common example of a tri-atomic alcohol. It has the composition (CH 2 )2 ' CH * (OH) 3 . CH 2 OH CHOH I CHsOH. It is so named on account of its sweet taste. It is a thick, clear syrup, miscible in water in all pro- portions and crystallizable in the absolutely pure state. Glycerine is a by-product of soap manufacture. CHAPTER XXX. SUGARS. THE compound C 6 Hi 4 is one of the paraffins, that is, it is composed of six methyl groups connected to form a chain in the following manner: H. H H H H Obviously several alcohol derivatives of this com- pound are possible. If one of the end H's were re- placed by OH the compound would be hexyl alcohol, C 6 Hi 3 OH. The other extreme, or complete hydroxy- lation, would of course consist 'of one OH group for every C atom. There could not be any more because if more than one OH group attaches itself to a C atom two of them combine and are split off as water (H 2 O). The compound resulting from the replacement of one of the H's on each C atom then is hexatomic alcohol : OH H H OH OH OH H C C C C C C H CC C C C C H OH OH H H H (Notice that the hydroxyl groups are not all on same side.) 166 SUGARS This compound is found in nature and is called mannite. It is a white crystalline substance with a sweet taste. It is fermentable and in other ways it is similar to sugars. Mannite is used extensively in bacteriology to differentiate certain kinds of bacilli, especially the organisms causing bacillary dysentery. Just one step, one slight chemical change, brings us to mannose, which is a sugar. By oxidizing one of the end groups we obtain an aldehyde. (Remember that formal- dehyde was produced by oxidizing methyl alcohol.) Mannose is a hexatomic alcohol (mannite) with an aldehyde H -c/o, : & group on the end: OH H H OH OH H _J_J '_'_'_ / -c c c c c- H OH OH H H If now we change the relative positions of the H and OH groups, we have the formula: OH H H OH H H I I I I I / TT f^ f~^ C^ (~* ^" 1 C^ C\ I ! I ! I H OH OH H OH This is grape sugar (glucose). Properties of Grape Sugar. Grape sugar is commonly known as glucose. (Glyc or glue stem in a word indi- cates sweetness, and the suffix ose shows that the sub- THE POLARISCOPE 167 stance is a sugar in the chemical sense.) Glucose is a white crystalline substance, soluble in water and slightly soluble in alcohol. It possesses a pleasant, sweet taste, but no odor. Glucose occurs in nature in combinations known as glucosides. It occurs in grapes uncombined, hence the name grape sugar. Commercially it is obtained by boiling starch with dilute acid (sulphuric acid is generally used because it can be so easily eliminated afterward). The reason why glucose can be obtained from starch in this manner will appear later. The Polariscope. A ray of light is said to be a vibration in ether. If we liken it to the vibration of a harp string we find that the excursions are not all in the same place. In other words, if the string were stretched in the direction exactly perpendicular to the earth the vibrations when the string is struck would not be confined to excursions north and south or east and west, but would swing to any or all points of the compass. Such is our conception of the vibrations of a ray of light. It is possible to keep the vibrations in the same plane: suppose two plane boards were placed one on each side of the string in such a manner that the string could move to and fro in one direction. The string would vibrate then in one plane. The same thing can be accomplished with light by allowing the ray to pass through a prism. The light waves vibrate in one plane just as the string did and the result is polarized light. If a ray of polarized light is passed through solutions of sugars, the vibration plane is turned to the right or left. Suppose the 168 SUGARS plane of vibration of a particular beam of polarized light is exactly vertical. The beam is passed through a solution of glucose. Now the plane of vibration is not vertical but has been rotated to the right (clock- wise) so many degrees. Glucose then is said to be dextro-rotary (dextra = right) . The amount of rotation is proportionate to the amount of sugar present. It is therefore possible to determine how much glucose is present in a solution without going through the process of recrystallizing it several times to purify it and finally drying and weighing it. The instrument with which one measures the rotation of the plane of light is called a polariscope. The number of degrees a prism must be rotated to bring the plane back to the original position is read and from this the amount of sugar can be calculated. The polariscope is used extensively in the sugar industries and in analyses of sugar-contain- ing substances. It is essential in accurate analysis of urines which contain sugar. The Asymmetric Carbon Atom. The power to rotate the plane of polarized light is due to the presence of a carbon atom which has all four of its bonds satisfied by different kinds of groups. Such a carbon atom is said to be asymmetric. If any two of these groups are the same, the carbon atom is not asymmetric. 1 Suppose, for example, that we substitute for three of the hydrogen atoms in methane the following groups (CHs), (OH), and (COOH), i. e., a methyl, a hydroxyl, and a carboxyl group. k One H atom remains. 1 Asymmetric means without, symmetry (not symmetrical). REDUCING POWER OF SUGARS 169 We have now a compound containing an asymmetric carbon atom, and according to the above statement should rotate the plane of polarized light. This compound H (OH) C COOH (CH 3 ) is lactic acid. It rotates the plane of polarized light to the right. Should we change the relative positions of the methyl and carboxyl groups we still have lactic acid, but this lactic acid rotates the plane to the left. This is the basis upon which our theories of optical activity of chemicals is built. Reducing Power of Sugars. Returning to the graphic formula of dextrose we find that it contains an aldehyde group, H I We remember that one of the chief characteristics of an aldehyde is its power to reduce substances, that is, take oxygen from other substances and become oxidized itself. To form an acid; the /^ Cf group becomes O X H || C OH. We therefore suspect dextrose of being able to reduce substances. Testing the substance for this property we find that it is so. If an alkaline solution of a copper 170 SUGARS salt be boiled in the presence of dextrose the copper is reduced and settles out as a bright red, fine precipitate (copper oxide, Cu 2 O). The solution best adapted for applying this fact in testing for dextrose is composed of copper tartrate and sodium hydroxide. This solution is usually known as Fehling's solution. Levulose. Another sugar, which on analysis would yield carbon, hydrogen and oxygen in exactly the same proportions as found in dextrose, rotates the plane of light to the left (contra clock-wise) and is, there- fore, called levulose. It is the sugar found in honey and also in fruits. The graphic formula of levulose differs from that of dextrose in that the CHO group is not at the end of the chain but inside. It is therefore not an aldehyde strictly but a ketone on account of the relative position of the aldehyde group. The strictly aldehyde sugars are called aldoses and ketone sugars named ketoses. Levulose reduces Fehling's solution like dextrose and can be fermented (i. e., broken up by ferments to form acids and carbon dioxide), though less readily than dextrose. Monosaccharids. Dextrose and levulose are typical examples of the several simple hexoses or monosac- charids. The group which they represent are called monosaccharids to distinguish them from the group, the members of which are formed by the chemical union of two simple sugars. SUMMARY OF CHAPTER XXX 171 SUMMAKY OF CHAPTER XXX. Hexane has the formula C 6 Hi4. It is possible to replace six H atoms with hydroxyl groups. No more OH groups can be joined to this compound because if more than one such group is attached to the same carbon atom, water would be split off by the union of the two. One hexa-hydroxy-hexane found in nature is man- nite. This alcohol is very similar in its physical proper- ties to sugars. It is used in bacteriological work. If one of the end alcohol groups of mannite be oxi- dized to an aldehyde group, the result is mannose, a sugar. By shifting the relative positions of the middle H and OH groups in mannose, grape sugar, another sugar is obtained. Glucose is a white crystalline substance, soluble in water and slightly soluble in alcohol, with a sweet agreeable taste and without odor. It is found in grapes and other fruits. It can be obtained by hydrolyzing starch. The polariscope is an instrument by means of which the rotation of the plane of polarized light is determined. In passing through sugar solutions the plane of polar- ized light is rotated to the right or left, and the amount of rotation is directly proportional to the amount of sugar present. This is a simple means for the deter- mination of sugars in solution. The power to rotate the plane of light depends upon the presence of an asymmetric carbon atom in the compound. On account of the presence of an aldehyde group 172 SUGARS in the hexoses (six carbon sugars), these compounds are capable of reducing substances like copper tartrate in alkaline solution. Both aldoses and ketoses (which see) have reducing power. The copper solution in alkaline tartrate used for testing sugars is commonly known as Fehling's solution. Levulose is fruit sugar. Solutions of this sugar rotate the plane of polarized light to the left or in opposite direction the rotation by grape sugar (dex- trose). Levulose is a ketose. It is found in honey. Dextrose and levulose are examples of simple sugars (hexoses) and are classed as monosaccharids in order to distinguish them from the di- and polysaccharids formed by the union of two or more monosaccharids. There are also sugars containing only five carbon atoms termed pentoses. They are sometimes found in human urine but are of more interest to the physio- logical chemist. CHAPTER XXXI. POLYSACCHARIDS. Cane Sugar. The familiar crystalline substance used extensively for sweetening is a result of chemical union between the two monosaccharids already de- scribed. Dextrose and levulose combined with the loss of one molecule of water, form cane sugar or saccharose. Both dextrose and levulose have the empirical formula CeHi 2 O 6 . Combination of these two produces Ci 2 H 24 Oi 2 , but there is lost in the process of uniting one molecule of water H 2 O, therefore we have as the empirical formula of saccharose Ci 2 H 22 On. Properties of Cane Sugar. Cane sugar is found in the juice of sugar cane, in beets, in bananas and other fruits. It crystallizes easily from concentrated solutions, and is therefore easily obtained in a pure state. Solu- tions of cane sugar rotate the plane of polarized light to the right. Fehling's solution is not reduced, show- ing that the aldehyde groups of the simple sugars composing it are completely masked. What evidence is there that cane sugar is composed of dextrose and levulose? Boil some cane sugar with weak hydrochloric acid (1 part concentrated HC1 to 100 c.c. solution of cane sugar) for two hours. The hydrochloric acid may be removed by precipitating I 174 POLYSACCHARIDS with lead or silver salts. Filtering through charcoal gives us a clear, colorless solution, which we may now compare with part of the original solution (cane sugar solution). The original solution rotated, the polar- ized light to the right, now the solution is levorotary (rotating to the left). The original solution did not reduce Fehling's solution but the new clear liquid does so very vigorously. From the original solution only one substance could be crystallized, from the recent liquid two substances can be separated, the one dex- trorotary the other levorotary. Further chemical test show these substances to be dextrose and levulose. It is thus proved that hydrochloric acid adds the lost molecule of water and splits cane sugar into its simple sugars. The word hydrolysis has been proposed to describe this process (hydro = water; lysis = breaking down). Thus this word is applied to all processes which add water chemically to a substance and divide it into its component parts. Invertase. We found that boiling with hydrochloric acid changed the power of the solution to rotate the plane of light, that is, it reversed the direction of rotation to almost an equal extent in the other di- rection. It practically inverted the rotation so we speak of the products of hydrolysis as invert sugar, and of the process as an inversion. Living yeast cells or extracts of yeast cells are also capable of bringing about the inversion of cane sugar, acting at the temperature of the human body. The substance in the yeast which acts in this manner is capable of converting many times its own weight of cane sugar, I OTHER DISACCHARIDS 175 it obeys certain laws of rate of reaction, and it is killed by heat. A substance answering such a description is known ,as a ferment. Names of ferments of this class end in ase, and the stem of the word indicates either their action or the substance they change. The logical name for the inverting ferment is invertase. Invertases occur in the intestinal juices having been secreted by the mucous lining of the intestine. It is by means of these ferments that the higher sugars (disaccharids) are reduced to simple sugars for absorp- tion in the process of digestion. Invertases are found in small amounts in the blood. OTHER DISACCHARIDS. Milk Sugar. If some copper sulphate is added to skim milk and then a small amount of sodium hydrox- ide, all the white substance of the milk will be co- agulated. Should we filter the mixture we obtain a clear liquid from which on evaporation a white, sweet substance crystallizes. This is milk sugar. The Latin word for milk is lactis, the ending -ose indicates a sugar; milk sugar is therefore well named lactose. Solutions of lactose rotate the plane of polarized light to the right, and also reduce Fehling's solution. Lactose is sweet, but not so sweet as glucose (dextrose), and it is not so soluble in water as the latter, nor is it so easily fermentable. Boiling with acid (hydrolysis) breaks up lactose into dextrose and another simple sugar, galactose. The mucous lining of the intestines yields a ferment known as lactase, capable of hydrolyzing lactose. 176 POLYSACCHARIDS In the digestion of lactose, this lactase hydrolyzes it before it is absorbed by the villi of the intestinal wall. Lactose is unique in being the only disaccharid of animal origin. It is found in the milk of all mammals : about 5 per cent, of cow's milk and about 7 per cent, of human milk is lactose. From whatever source it is obtained it is the same chemically. Malt Sugar. Malt sugar is a disaccharid composed of two molecules of dextrose, yielding these on hydro- lysis by boiling with an acid or under the influence of maltase. Maltose is sweet, is dextrorotary, and reduces Fehling's solution. It is found in malt where it is produced from starch by the action of a ferment derived from germinating (sprouting) grain. STARCHES. The reserve material of plants is stored in the form of pure white insoluble substances called starches. Most of the dry material of the potato, for example, is starch; corn also contains a high percentage of this carbohydrate. A prominent characteristic of starch is its reaction with iodine to form a blue color, which disappears on heating but reappears on cooling. To test any mixture for the presence of either of these substances (starch or iodine) one has merely to add the other. Under the microscope starch is found to be com- posed of small, oval grains, having on them lines forming elongated ellipses drawn about a common centre. Starches from different sources have different SUMMARY OF CHAPTER XXXI 177 shapes and markings, but so far as is known they are the same chemically. In fact the exact composition, of starch is not known. It is evident from many studies that it is composed of a group of molecules of dextrose, but how many and their manner of combination has not been determined. Boiling with mineral acids decomposes starch with the formation of dextrose (glucose) and dextrins. A ferment, amylase (amylum = starch), found in the saliva, in pancreatic juice and in sprouting grain is also capable of hydrolyzing starch in the same manner. The starch grains have an outer coat of cellulose (wood tissue) which protects it from ferment action until it .is broken. Heat ruptures this coat so that cooking prepares the starch for digestion. In this we find the reason for the indigestibility of raw starch, except by cows and goats, whose intestinal juices contain ferments capable of dissolving cellulose. Starch is the great source of commercial glucose. Grain alcohol is made by hydrolysis and fermentation of starches. (See Ethyl Alcohol.) Cellulose. Cellulose is a complex compound of sugars and starches in chemical combination. It is characterized by its extreme insolubility. Wood fibre is largely cellulose. SUMMARY OF CHAPTER XXXI. Polysaccharids are carbohydrates made up of two or more monosaccharids. Disaccharids are polysac- charids consisting of two monosaccharids, 12 178 POLYSACCHARIDS Cane sugar is a disaccharid composed of dextrose and levulose; that is, 1 molecule of cane sugar is composed of 1 molecule of dextrose, plus 1 molecule levulose, minus 1 molecule water. Dextrose (C 6 Hi 2 O 6 ) + levulose (C C=O + H 2 O. I I O.NH 4 NH 2 Urea. Urea is decomposed by HN0 2 (nitrous acid) into gaseous nitrogen and CO 2 . The CO 2 can be absorbed by NaOH and the amount of N gas measured. The amount of urea can then be calculated. (See chapter on Uranalysis.) It has been stated (see proteins) that proteins are broken down during digestion into amino-acids. The amino-acids are absorbed and at some unknown place are reassembled to form proteins when they are neces- sary for growth or repair. Excess of proteins are not stored in the body as fat and sugars are. The excess of amino-acids are oxidized in the liver to form urea and ORGANIC CONSTITUENTS 245 excreted in the urine. Therefore the amount of urea in the urine varies with the amount of proteins ingested. Urea is also excreted in the sweat. Uric ^4cid. Although the actual amount of uric acid excreted in twenty-four hours is comparatively small (about 0.5 gram) it is considered an important constituent of the urine. It is supposed to come from the breaking down of the nuclei of body cells and from the nuclei of cells in the meat ingested. The crystals of uric acid may assume several forms: like wedges, prisms, plates, or dumb bells. Free uric acid is insoluble in water but its lithium and its sodium salts are easily soluble. Ammonium urate is very slightly soluble. Crystals of uric acid or ammonium urate are colorless when pure, but they occlude coloring matters so that often the brownish-red precipitate seen in urines, after standing, consists of uric crystals or its salts. Uric acid in excess reduces Fehling's solution slightly, and may, therefore, when present in large amounts, interfere with sugar tests in the urine. The formula for uric acid is: HN O=C - C NH CO. HN C NH It is a tri-oxy-purine. On oxidation the three C atoms connected with one another are changed to CO 2 and two molecules of urea are given off. Indican. Indoxyl potassium sulphate or indican is one of the most important of the ethereal sulphuric 246 THE URINE acids found in the urine. It is oxidized by chlorine (bleaching powder) to indican blue or indican red. (See test for Indican in Appendix.) Indican is a product of putrefaction and when found in the urine is evidence that putrefaction is taking place in the intestinal contents. Creatinin seems to be a product of cell activity, and on a creatin-free diet the amount excreted seems to be fixed for each individual. Creatin occurs in muscle: creatinin is reduced creatin. For formula and prop- erties, refer to any text-book on physiological chemistry. Other Organic Constituents of the Urine. A greaf. many other organic compounds of more or less com- plexity may occur in variable quantities. Their presence in small quantities may be normal, but the extent to which they may be present to constitute abnormality varies under different conditions. SUMMARY OF CHAPTER XL. The amount of urine excreted varies markedly. The temperature and humidity, the amount of water drunk, the kind and quantity of food ingested, and the blood-pressure affect the secretion of urine. The average varies from 800 to 1200 c.c. per day. The specific gravity varies with the volume and with diseased conditions. It generally varies from 1015 to 1025, though it may exceed these limits under certain conditions and still be normal. Under normal conditions, urine of high specific gravity makes us SUMMARY OF CHAPTER XL 247 suspicious of diabetes and tests for sugar in urine are accordingly made. The color is also subject to considerable variation from pale yellow to deep orange. The normal color is due to urochrome. Bile salts are found in abnormal conditions. Substances taken into the stomach may be eliminated in the urine and change its color. The odor of the urine is said to be due to a substance of unknown composition called uronoid. Substances in the food may impart a characteristic odor to the urine, e. g., asparagus. The reaction of fresh, normal, human urine is slightly acid, although after the ingestion of large amounts of vegetables or organic acids it may be alkaline. De- composition of various nitrogenous bodies in the urine by ferment action soon renders the urine alkaline. This alkalinity results from the ammonia liberated. The salts of the urine consist mainly of NaCl. About 12 grams of NaCl per day are excreted. Chlorides of K, (NH 4 ), Mg, and Ca are also found. Phosphates of alkaline metals and alkaline earths occur. The three types of phosphates found may be represented as follows: (1) M 3 PO 4 ; (2) M 2 HPO 4 ; and (3) MH 2 PO 4 , in which M represents any monovalent metal like Na or K. Sulphates (inorganic and organic) and small amounts of carbonates and nitrates are found. The most important organic constituent is urea. 90 per cent, of the total N is urea N. 30 to 40 grams of urea are excreted by an average man in one day. Urea can be separated from urine by adding HNO 3 . The urea nitrate can be recrystallized and finally decomposed by BaCO 3 . Urea has the composition 248 THE URINE CO(NH 2 ) 2 . It can be made synthetically by heat- ing ammonium carbonate. Ammonium carbamate CO.NH 2 .O.NH 4 is first formed and later changes to urea on losing 1 molecule H 2 O. Urea is formed in the liver by the oxidation of the groups of amino-acids which have not been utilized to rebuild proteins. Urea is decomposed by HN0 2 forming gaseous nitrogen and C0 2 . Uric acid is supposed to result from the breaking down of the nuclei of ingested cell nuclei. On standing, urine deposits crystals of urine acid, especially after the urine becomes alkaline. Uric acid forms salts with Na, K, and (NH 4 ) groups. Sodium and potassium urates are much more soluble in water than ammonium urate. Uric acid in excess reduces Fehling's solution and may thus be a source of error in examining a sample of urine for sugar. Uric acid is a tri-oxy-purine (closely related to the alkaloid caffeine). Urea is formed by the oxidation of uric acid. Indican is indoxyl potassium sulphate. When this body is found in urine it indicates intestinal putrefac- tion. It may be detected in urine containing it by treating with bleaching powder. A blue or red color indicates the presence of indican. Creatinin seems to be a product of cell activity, and on a creatin-free diet the urine of individuals contains a fairly constant amount of creatinin peculiar to that person. Creatin occurs in muscle: creatinin is reduced creatin. Other organic compounds may be found at times in the urine, but are of interest more to the students of physiological and pathological chemistry. CHAPTER XL1. URANALYSIS. THE normal appearance of a sample of urine is no guarantee that it is normal. It is very important that specimens should be tested, even though they appear to be perfectly normal. Collection of the Urine. For chemical analysis of urine it is extremely important that a well-mixed twenty-four-hour specimen be collected, for the reason that the character of the various voidings are subject to wide variations. Sometimes it is desired that the night urine be kept separate. In this case the voidings between 9 P.M. and 6 A.M. are poured together. A twenty-four-hour specimen includes voidings beginning with the first after 6 A.M. and including the 6 A.M. voiding on the next day. It is very essential that all voidings should be saved and the total accurately measured. Preservation of Urine. All specimens should be kept in the ice-box until immediately before examination. Urine is easily fermented the reaction changes, urea is decomposed, the ammonia content increases, carbohydrates may be fermented, and the microscopic elements decompose if the specimen is not kept cold. Various preservatives have been recommended for 250 . URANALYSIS special kinds of work. The analyst should know the preservation used. For chemical work chloroform is perhaps the best. Add 5 c.c. for each liter of urine. The urine should be well shaken after each addition and the bottle kept tightly corked. The chloroform settles to the bottom and adds nothing to the volume. Immediately before analysis the urine may be poured off the chloroform and a sample exposed to incubator temperature for a very few minutes, or air bubbled through in order to remove traces of chloroform. Any larger trace of chloroform may yield a suggestive reducing reaction with Fehling's solution. 1 If formalin is used one is more liable to obtain suggestive sugar reactions with Fehling's solution, yet for microscopic examination formalin seems the best preservative. Thymol is used sometimes, but a test similar to bile may be obtained in the specimen. Amount of Urine. The amount of urine should be carefully measured in a cylinder registering volume in cubic centimeters. Specific Gravity. For clinical examination the specific gravity is approximated by using the urinometer. The instrument is standardized to read 1000 in distilled water at 15 C. The glass cylinder should be clean and dry. It is filled about four-fifths full of urine by gently pouring in order to prevent foaming. Foam should be re- moved with a strip of filter paper. 1 Formalin does not reduce Benedict's solution but chloroform does. For each liter of urine 2.5 c.c. formalin is sufficient. COLOR 251 The temperature is important. A small thermometer should be placed in the urine and stirred a few times. If the temperature is not 15 C. the jar should be placed in warm or cool water, as the case requires, and gently stirred until the thermometer registers 15. If the proper temperature cannot be attained in a hurried examination, a correction by adding for every 3 above 15 is made. For example, a urine reading 1022 at 21 should be corrected to read 1024. The bobbin should be clean and dry. It is inserted carefully and gently tapped and read when it comes to rest. It is important that the bobbin does not touch the side or bottom of the cylinder. The graduation on level with the surface of the urine (the bottom of the meniscus) is read and recorded. In a child's or catheterized specimen there may not be sufficient urine to fill the cylinder. Dilute with equal parts of water and test as above. Multiply the last two figures by 2 and add to 1000. If reading is 1010, then 10 X 2 = 20; 20 + 1000 = 1020 (correct reading) . Total Solids. When requested, estimate total soMds by multiplying the last two figures of the specific gravity by Haser's empirical coefficient 2.33. This will give grams per liter. To find percentage divide by 10. Color. The color of a given depth of urine viewed against a white background is recorded in well-recog- nized terms : straw, light yellow, yellow, amber, orange, brown, brownish red, etc. If the foam is yellow, bile is present. The foam is otherwise white. 252 UR ANALYSIS The character of any precipitate is described as to amount (small, moderate, abundant), color, and con- sistency (dense or flocculent). The general appearance of the urine (clear, cloudy) is observed. Odor. A small amount of urine is brought to boiling in a beaker covered with a watch-glass. The glass is removed and odor recorded as fruity, ammoniacal, sulphidic, etc. Reaction. A strip of blue and of red litmus paper are immersed to half their length in the urine and immediately withdrawn. The color change is observed. If the blue paper changes quickly to red, the urine is strongly acid; a slight change is interpreted as faintly acid. The strips are laid on white, clean, dry filter paper. If the red paper changes blue, the urine is recorded as being alkaline. If the strip of litmus is allowed to dry and becomes red again, the alkalinity is recorded as being due to ammonia, otherwise it is due to a fixed alkali. Normal urine is always faintly acid, varying with the diet. Total Acidity. When required, the total acidity may be determined by titrating with yV NaOH, accord- ing to the method of Folin. 25 c.c. of urine are delivered from a pipette into a 200 c.c. Erlenmeyer flask. 100 c.c. distilled water, 20 grams potassium oxalate, and 2 drops of a 0.5 per cent, alcoholic solution of phenolphthalein are added. The flask is shaken well for one minute and -$ NaOH added from a burette until a distinct reddish color is produced. The flask is shaken after each addition. The amount of NaOH necessary to produce color when REDUCING SUGARS 253 multiplied by 4 gives the number of c.c. $ alkali required per 100 c.c. urine. Albumin. Albumin is coagulated and precipitated by heat. Pour into a clean chemical test-tube 2 c.c. of a saturated solution of NaCl in water. Pour in filtered, clear, slightly acid urine until three- quarters full. The top zone of f to 1 inch is heated to boiling in a small flame (alcohol lamp preferred). If a whitish cloud is observed in the heated zone, when viewed against a black background, add one drop of 25 per cent, acetic acid and boil again. Repeat the addition of acetic acid and the boiling three times. If the precipitate is due to phosphates this treatment will cause it to disappear. If the cloudiness is due to albumin it will persist and perhaps increase. The whitish deposit sometimes formed on the glass by the flame, especially when gas is used, should not be con- fused with an albumin test. If much albumin is present a heavy flocculent precipitate will be formed. Heller's Nitric Acid Test. Place 10 c.c. of clear, concentrated nitric acid in a two-ounce conical stand glass. (If the nitric acid is yellow, boil in a beaker until clear.) With a pipette carefully overlay about 25 c.c. filtered urine over the nitric acid. After three minutes observe the line of contact against a dark background. A precipitate at the juncture of the two liquids indicates albumin. The red or reddish-violet ring often observed should not be mistaken for a precipitate. Reducing Sugars. Fehling's Method. Solution A. Dissolve 34.65 grams powdered CuSCX in 300 c.c. warm, distilled water in a 500 c.c. volumetric flask. 254 UR ANALYSIS Allow to cool and make up to mark (500 c.c.) with cold, distilled water. Solution B. Weigh out roughly on filter paper 125 grams potassium hydroxide. (KOH is caustic; handle with forceps. Break sticks in a large mortar.) Place in a 500 c.c. beaker and add 300 c.c. distilled \vater. When all is dissolved pour into a 500 c.c. volumetric flask and rinse three times with small por- tions of distilled water, pouring each portion into the flask. Pulverize about 180 grams sodium-potassium tartrate (Rochelle salt) in a mortar. Weigh out exactly on glazed paper 173 grams of powdered Rochelle salt and pour carefully into flask containing the KOH solution. Shake until all is dissolved; allow to cool and make up to 500 c.c. mark with distilled water; preserve in a rubber-stoppered bottle. For tests, mix equal parts of solutions A and B immediately before using. To 1 c.c. of the mixture in a test-tube add 4 c.c. distilled water and boil. (If a precipitate is formed the solutions are worthless.) To the warm solution add three or four drops of the urine to be tested and boil. Repeat this several times. A yellow or brownish precipitate indicates the presence of reducing sugars. Phosphates and uric acid may form precipitates which may be confused with the positive reaction. Fehling's solution should not be used after it has stood over long periods. Benedict's Method. Dissolve 173 grams sodium citrate and 100 grams anhydrous sodium carbonate in 600 c.c. distilled water. Pour through a folded filter into a glass graduate and make up to 850 c.c. with PROCEDURE 255 water. Dissolve 17.3 grams powdered copper sulphate in 100 c.c. water and make up to 150 c.c. with water. Place solution 1 (carbonate-citrate solution) in a large beaker and add the copper solution slowly with constant stir- ring. The mixed solution is stored in a rubber-stoppered bottle. It does not deteriorate on long standing. Procedure. To 500 of Benedict's solution, in a test-tube, add eight drops of the urine to be tested and boil for three minutes. Allow to cool spontane- ously (do not cool under tap). If glucose is present a large precipitate will form, otherwise the solution will remain perfectly clear or only slightly turbid. Quantitative Estimation of Dextrose. Benedict's Method. Benedict's quantitative solution is made up as follows: Dissolve in 750 c.c. hot distilled water the following: crystals of sodium carbonate, 200 grams (or anhydrous 100); sodium or potassium citrate 200 grams, and 125 grains potassium sulphocyanate. When all is dissolved filter if the solution is not clear. Dissolve exactly 18 grams pulverized, pure crystals of copper sulphate in 100 c.c. water and pour slowly into solution No. 1, with constant stirring. Allow to cool, add 5 c.c. of a 5 per cent, aqueous solution of potassium ferrocyanide solution and make up to 1000 c.c. with distilled water. This solution keeps indefinitely in a rubber-stoppered bottle. Procedure. If the urine 1 to be titrated shows very heavy reduction (i. e., contains probably a large amount 1 In case chloroform was used for preservation, a portion of the urine should be brought to boiling-point and quickly cooled. This rids it of the chloroform. 256 URANALYSIS of sugar) dilute 10 c.c. of it to 100 c.c. with distilled water. Fill a clean and dry 50 c.c. burette with the solution. Exactly 25 c.c. of Benedict's solution is measured with a pipette into a porcelain evaporating dish, 25 cm. in diameter, and to it about 15 grams of crystal- lized sodium carbonate (or 8 grams dry Na 2 COs) are added. Half a teaspoonful of powdered pumice or talc is added and the dish heated to boiling over a free flame until the carbonate has entirely dissolved. The diluted urine is now run from the burette rather rapidly until a chalk-white precipitate forms, and the blue color of the mixture begins to lessen perceptibly. Now the diluted urine must be dropped from the burette very slowly a few drops at a time until the disappearance of the last trace of blue color which marks the end-point. The solution must be kept boiling vigorously throughout the entire titration. If the boiling mixture begins to bump or spatter add a little distilled water to make up for the loss by evaporation. Calculation. The amount of diluted urine which was necessary to reach the end-point is found by reading the burette. This figure divided by 10 gives the equivalent amount of urine (because the burette con- tained urine diluted ten times). This amount of urine must contain 50 mg. (0.05 gram) of dextrose, because it requires 50 mg. dextrose to reduce 25 c.c. Benedict's solution. Let x = amount of urine containing 50 mg. dextrose, then ^ x 100 = per cent, sugar in original urine. UREASE METHOD 257 Urea. Hypobromite Method. 1 Solution A. Dissolve 62.5 grams sodium bromide in 400 c.c. water. Pour into a 500 c.c. volumetric flask. Add 22 c.c. pure bromine (under hood bromine fumes are very irritat- ing), and make up to 500 c.c. with distilled water. Preserve in rubber-stoppered bottles. Solution B. Dissolve 125 grams sodium hydroxide in 400 c.c. distilled water. Allow to cool. Pour into a volumetric flask and make up to 500 c.c. Preserve in a rubber-stoppered bottle. Into the open arm of the clean Doremus-Hinds ureometer add a drop or two of urine; turn the stop- cock so that the urine just fills the lumen and turn to off position. Fill the closed arm with a mixture of equal parts of solutions A and B, and tilt to let out any air. Set upright and fill open arm with urine, avoiding bubbles of air. Turn stop-cock and allow 1 c.c. urine to flow in. After twenty minutes read the amount of gas in the closed arm. Each small sub- division of gas represents 0.001 gram of urea per cubic centimeter of urine. If percentage is desired move the decimal point two places to the right. Every small division means 0.1 per cent. urea. Urease Method. Van Slyke's Modification of Mar- shall's Method. This method is accurate and so simple that the nurse with little laboratory experience can be 1 The following method for the preparation of the hypobromite solution has also been recommended. Keep on hand in a rubber- stoppered bottle a 20 per cent, solution of NaOH. When ready for a test add to 40 c.c. of this solution 1 c.c. of pure bromine and shake. This solution may be used in place of the mixture of A and B, as indicated in the text. 17 258 UR ANALYSIS taught to use it. The determinations, especially the first, should be done under the direction of a physio- logical chemist. The details of the method are given for reference of the nurse after she has learned to apply it. The method depends upon the fact that urease, a ferment found in the soy bean, is capable of convert- ing urea into ammonia without any loss of nitrogen. The ammonia can then be blown out of the urine into a known amount of standard acid. The acid is titrated at the end of the experiment to determine how much has been neutralized by the ammonia. The method is rapid and of advantage on account of the fact that several analyses may be run at the same time. PROCEDURE. One-half c.c. of urine 1 is measured into the bottom of tube A. Exactly 5 c.c. of a solution containing 6 grams of KH 2 P0 4 per liter are then run in from a burette, and 1 c.c., accurately measured, of a 10 per cent, solution of urease 2 is added. 1 An Ostwald pipette is used, the stem of which is a heavy walled capillary tube of only 1 mm. bore. The pipette, which should deliver in about twenty seconds, is calibrated by weight for blow-out delivery, and permits measurement with an accuracy of 0.001 c.c. The pipette is allowed to deliver with its tip against the lower part of the test-tube wall until the bulb is emptied; the remainder of the contents is then blown out. These pipettes, as well as the 100 c.c. test-tubes of special heavy glass, provided with inlet and outlet for aeration, the block holder shown in the figure, and a brass aspirator pump suitable for the method can be obtained from Emil Greiner, 45 Cliff Street, New York. 2 The enzyme preparation used should be standardized as follows: A solution is made containing 3 grams, accurately weighed, of pure urea per 100 c.c. Using the special pipette described in the urine analysis, one measures into tube A, 0.5 c.c. of the above urea solution, 5 c.c. of 0.6 per cent. KH2PO4, and the amount of enzyme solution intended to be used in analysis (usually 1 c.c. of 10 per cent, enzyme). The reaction is allowed to run at room temperature (or 50 if desired) UREASE METHOD 259 The solutions in the tube are well mixed, 2 drops of caprylic alcohol to prevent subsequent foaming are added, and the stopper bearing the aerating tubes shown in the figure is put into place. Twenty minutes at a room temperature of 15, or fifteen minutes at 20 or above, are allowed for complete decomposition of urea. No harm is done if the solutions are allowed to stand longer, but the time must not be cut shorter unless more enzyme is used. While the enzyme is acting, one measures 25 c.c. of -^ hydrochloric or sulphuric acid into tube B and connects the tubes as shown in the figure. After the time for complete decomposition of urea has elapsed the air current is passed for a half minute in order to sweep over into B a small amount of ammonia which has escaped into the air space of A during the decomposition. A is now opened and 4 to 5 grams of potassium carbonate measured roughly from a spoon are poured in (in order to assure most rapid removal of ammonia by air current it is necessary to have the solution at least half saturated with carbonate). The air current is now passed rapidly through the tubes until all the ammonia has been driven over into the acid in B. The time required for this depends on the speed of for the length of time allowed in analysis, and the ammonia is deter- mined as described for urine analyses. It should neutralize 25 c.c. of -^ acid. If it falls slightly short, it is well to repeat the test, doubling the time interval, as some samples of urea are not 100 per cent, pure, and the short figure may be the fault of the urea, not of the enzyme. If in the longer interval no more ammonia is formed than in the shorter, the urea decomposition was complete in the shorter time, and the enzyme is sufficiently active. This method is described in detail in the article by Van Slyke and Cullen, Journal Biological Chemistry, vol. xix, No. 2, October, 1914. 260 UR ANALYSIS the air current. With a rapid pump or house vacuum it is possible to aerate completely in five minutes; Tb Pump -Wash Bottle Apparatus for determining urea content by means of urease. (Journal of Biological Chemistry.) while a slow pump may require a half hour. The time required for complete aeration is determined for the BILE 261 particular vacuum used by trial, and a safe margin allowed in the determinations. When the aeration is finished the excess acid in B is titrated with ^ NaOH. The operations can be concisely summarized in the following diagrammatic form: 0.5 c.c. urine. , ,, . , 5.0 c.c. 0.6 per cent. KH 2 P04. 1. Measure into A . 1Q ^ 1Q p P er ^ urease> 2 drops caprylic alcohol. Place stopper as shown in Fig. 1 and let stand fifteen minutes. {25 c.c. 5 n c acid. 1 tUZ^i^. 8 * 1 " 1 drop caprylic alcohol. 3. (After 15 minutes standing) aerate one-half minute. Then open A and add 4 to 5 grams K^COs. 4. Aerate all NHs from A over into B. 5. Titrate excess acid in B with ^ NaOH. 6. Calculate: 0.056 X c.c. J^ acid = grams urea + ammonia nitrogen per 100 c.c. urine. In order to determine the ammonia nitrogen alone one measures 5 c.c. of urine into A, adds the carbonate at once, and aerates as described above. The acid neutralized is multiplied in this case by the factor 0.0056, to give the per cent, of ammonia nitrogen. No extra time is required for the ammonia determina- tion, as one merely aerates the extra pair of tubes in series with the same air current used for the ammonia + urea determination. As a matter of fact, one can conveniently run as many as eight pairs of tubes on the same air current, taking the precaution at the end of the aeration to disconnect the series in the middle first in order to prevent back suction. Bile. Gmelin's Test. The urine is superimposed over nitric acid in exactly the same way as in the 262 URANALYSIS test for protein. In the bile test, however, the nitric acid should be slightly yellow instead of clear. Nitric acid may be turned yellow by adding a few pine shavings (Emerson). The yellow color indicates the presence of nitrous acid, HNO 2 . In the presence of bile the line of contact of the two liquids will present strata of colors varying from green, blue, violet, red, and yellow. In urine containing bile, the foam is always yellow otherwise it is white. Indican. Jaffe's Test. To 5 c.c. urine in a test- tube add 5 c.c. concentrated HC1. To the mixture add 2 c.c. chloroform and 5 drops of a filtered saturated solution of bleaching powder. Shake thoroughly. A blue or red coloration of the chloroform indicates that indican was present in the urine. Diazo Reaction. Solution A. 5 grams of sodium nitrite in 1 liter water (this solution should not be used after standing two weeks). Preserve in glass-stoppered bottle. Solution B. 5 grams of sulphanilic acid and 50 c.c. HC1 in 1 liter distilled water. Preserve in glass-stop- pered bottle. Mix solutions A 1 part to 50 parts solution B. To 5 c.c. urine add 5 c.c. mixture of reagents as above indicated. Mix thoroughly by shaking, add quickly 1 c.c. ammonium hydroxide. If the fluid and the foam turn red the test is positive. On standing a precipitate is formed leaving a supernatant fluid, which is green, blue, or violet. In normal urine the reagent produces a brownish-yellow color. INDEX. A ACETALDEHYDE, 157 Acetanilid, 210 Acetic acid, 158 Acetylene, 121. See Calcium carbide. Acids, 106 acetic, 158, 183 amino-, 204 carbolic, 195 fatty, 185 formic, 146, 183 hydrochloric, 83 lactic, 169 organic, 146 propionic, 183 unsaturated, 185 uric, 245 Albumin in urine, 253 Albumins, 219 Alcohol, butyl, 161 di-atomic, 162 ethyl, 154 grain, 154 higher, 161 methyl, 144 mon-atomic, 162 primary, 161 propyl, 161 secondary, 162 tri-atomic, 162 wood, 154 Aldose, 170 Alkaloids, 212 Alum, 129 Aluminum, 129 Amalgam, 125 Amino-acid, 204 Ammonia, 202 Ammoniated mercury, 127 Ammonium acetate, 203 hydroxide, 203 Analysis, volumetric, 109 Anilin, 209 Animal sugar, 175 Anode, 43, 48 Antimony, 115 Antitoxin, 228 Argyrpl, 134 Arsenic, 114 antidote, 130 Arsine, 114 Asymmetric carbon atom, 168 Atomic weights, 39 Atoms, 30 Avogadro's hypothesis, 40 B BACTERIA, nitrogen-fixing, 202 Baking powder, 103 Bases, 106 Benedict's method, 254, 255 Benzaldehyde, 198 Benzene, 191, 192 ring, 193 series, 191 Benzol. See Benzene. Bichloride of mercury, 126 Bile, 261 test for, 262 Bismuth, 116 Bitter almonds, oil of, 199 Bleaching powder, 86, 120 264 INDEX Blood, 226 alkalinity of, 227 amount of, 230 cells, 228 clot, 226 functions of, 231 salts of, 227 serum, 227 specific gravity of, 230 Boiling-point, 57 Borax, 69 Bromine, 89 Bromoform, 143 Bunsen burner, 121 Butane, 160 CALCIUM, 119 carbide, 139 Cane sugar, 173 Carbohydrates, 180 Carbon, 138 monoxide, 138 Carbonates, 138 Carboxyl group, 146, 168, 169 Casein, 235 Catalyzers, 78, 135 Cathode, 43, 48 CeUulose, 177 Centigrade, 56 Charcoal, 138 Chemical change, 23 Chloral, 157 Chlorine, 81 preparation of, 81 properties of, 82 uses of, 82 Chlormethane, 143 Chloroform, 143 Coal tar, 191 dyes, 210 Colloids, 135, 227 Condensation, 60 Conservation of mass, 32 Corrosive sublimate, 126 Creatinin, 246 Crenation, 230 Cresols, 198 Crystallization, 66 Cyanogen, 208 DEAD Sea, 81 Destruction of matter, 31 Developer, 134 Dextrose, 168, 173 in blood, 231 Dialysis, 135, 219, 227 Diamond, 138 Diazo reaction, 262 Diazonium, 211 Dietary, carbohydrates in, 182 proteins in, 222 Digestion, carbohydrate, 180 intestinal, 180 mouth, 180 pancreatic, 180 stomach, 180 Disaccharid, 175 Distillation, 66 fractional, 67 E ELEMENTS, 27 Empirical formula, 142 Energy, 32, 181 kinetic, 33 latent, 33 Epsom salt, 123, 124 Ethers, 150 Evaporation, 60 FAHRENHEIT, 56 Fats, 183 body, 185 digestion of, 187 milk, 236 vegetable, 186 Fehling's solution, 170, 173, 253 Fermentation, 156 Ferments, 156, 159 in blood, 231 oxydizing, 231 Fibrin, 226, 232 Filter, porcelain, 70 sand, 70 INDEX 265 Flame test, 121, 123 Fluorine, 93 Fluorite, 119 Food value of carbohydrate, 182 of fats, 187 of milk, 237 of proteins, 221 Fool's gold, 95, 114 Formaldehyde, 144 Formalin, 144 Formic acid, 146 Formula, graphic, 169 Fowler's solution, 115 Fractional distillation, 67 Freezing, 61 GASOLINE, 161 Globulins, 218 Glucose, 167, 173 Glycerine, 162, 184 Glycogen, 181 Glycoproteids, 220 Gmelin's test, 261 Gram, 24 Grape sugar, 166 Graphite, 138 HARDNESS of water, permanent, 69 temporary, 69 Heat, 60 of condensation, 60 of evaporation, 60 of freezing, 62 of melting, 62 Heller's test, 253 Hemolysis, 230 Hexane, 171 Hexose, 170 Homogenized milk, 236 Honey, 170 Hydrated lime, 21 Hydrocarbons, saturated, 160, 163 Hydrogen, 48 Hydrogen, nascent, 49 peroxide, 77 properties of, 49 uses of, 49 Hydrolysis, 174 Hydrometer, 55 Hydroxyl, 73 Hygroscopic, 121, 162 "Hypo," 134 Hypobromite solution, 257 INCOMPATIBLES, 75 Indican, 245 test for, 262 Indicators, 108 Indigo, 209, 210 Inversion, 174 Invert sugar, 174 Invertase, 174 Iodine, 90 preparation of, 91 properties of, 92 lodoform, 143 lonization, 73 Iron, 129 Isotonic, 229 JAFFE'S test, 262 KETOSE, 170 Kjeldahl, 225 LACTASE, 175 Lactose, 175, 236 Laking of blood, 230 Law of combining weights, 36 of conservation of mass, 32 of constant proportions, 35 of multiple proportions, 35 Lead, 133 266 INDEX Levulose, 170, 173 Light, polarized, 167 Lime, 19 Lipase, 187 Liquor chlori compositus, 82 Lithium, 110 Litmus, 85, 108 Lunar caustic, 133 Lysis, 174 M MAGNESIA, 124 Magnesium, 123 Malt sugar, 156, 175 Manganese, 131 Mannite, 166 Marsh gas, 141 series, 153 test, 115 Marshall's method, 257 Melting, 62 Mercury, 125 Meta compounds, 195 Metals, 25 Methane, 142 Methyl acohol, 144 Milk, 235 clot, 235 cow's, 237 human, 237 salts, 236 sugar, 175, 236 Mine, damp, 141 Moisson, 138 Molecular solutions, 106 Molecules, 30 Monosaccharid, 170 N NITRIC acid test for albumin' 253 Nitro benzene, 193 Nitrogen, 202 OIL of bitter almonds, 199 corn, 186 Oil, cottonseed, 186 olive, 186 peanut, 186 Olein, 185 Ortho compounds, 195 Osmosis, 228 Oxidation, 27 Oxides, 45 Oxygen, 26, 43 preparation of, 43 properties of, 44 uses of, 44 Ozone, 46 PALMITIN, 185 Para compounds, 195 Paraffins, 160 Paraldehyde, 157 Pentane, 160 Pentoses, 172 Petroleum, 160 Phenol, 195 Phenylhydrazin, 210 Phosphorus, 112 effects of, 113 forms of, 113 uses of, 113 Phot9graphy, 134 Physical change, 22 Plasma, 226 Plaster of Paris, 120 Platinum, 134 Polariscope, 167 Polysaccharid, 173 Potassium, 109 permanganate, 131 sulphocyanate, 209 Pressure, osmotic, 229 Pyridin, 213 Priestly, 27 Propane, 158, 160 Proportions, constant, 35 multiple, 35 Proteids, 218 compound, 219 ' derived, 220 glycoproteids, 220 simnle, 218 Proteins, 216 INDEX 267 Proteins, classification of, 218 digestion of, 220 occurrence of, 217 Pro-thrombin, 232 Purification of substances, 66 Putrefaction, 156 QUICK-LIME, 19 R RADICAL, organic, 144 Reducing power, 169 Reversible reaction, 51, 119 Rusting process, 25 SACCHAROSE, 173 Saline, 101, 230 Salt, 102 Saltpetre, 110 Salvarsan, 115 Silver, 133 Soda, 103 lye, 104 Sodium, 100 chloride, 102 hydroxide, 104 Solubility, effect of tempera- ture on, 65 Solutions, 64 Fehling's, 170, 173 hypertonic, 229 hypotonic, 229 Specific gravity, 55 of urine, 240, 250 Spectroscope, 124 Spectrum, 124 Springs, 68 Starch, 176 Stearin, 185 Strontium, 121 Structural formula, 142 Sugars, 165 of lead, 133 Sugars, reducing of, in urine, 253 Benedict's method, 254 Fehling's method, 253 Sulphocyanates, 209 Sulphur, 95 dioxide, 97 Symbols, 28 TARTAR emetic, 116 Thermometer, 56 Thrombin, 232 Toluene, 197 Toluol. See Toluene. Tonicity, 229 URANALYSIS, 249 Urea, 243 constitution of, 244 estimation of, 257 hypobromite method, 257 urease method, 257 Van Slyke's method, 257 Urease method, 257 Ureometer, 257 Uric acid, 245 Urine, 240 acidity of, 252 albumin in, 253 amount of, 240, 250 collection of, 249 color of, 241, 251 odor of, 241, 252 organic bodies in, 243 preservation of, 249 reaction of, 242, 252 salts in, 242 specific gravity of, 240, 250 total solids in, 251 Urochrome, 241 Uronoid, 241 VALENCE, 97 Van Slyke method, 257 268 INDEX Vinegar, 158 method of, 158 Vitamines, 213 Volumetric analysis, 109 W WATER, 53 composition of, 72 hard, 69 potable, 70 properties of, 54 rain, 70 river, 68 sea, 68 soft, 69 Water, as standard of compari- son, 55 synthesis of, 72 uses of, 54 Weights, 24 Wood, 177 XANTHOPROTEIC reaction, 218 X-rays, 117 Xylene, 197 Xylol. See Xylene. 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