LIBRARY OF THE UNIVERSITY OF CALIFORNIA. GIFT OK Accession 2g.43.69..... Class ... TWENTIETH CENTURY TEXT-BOOKS EDITED BY A. F. NIGHTINGALE, PH. D. SUPERINTENDENT OF HIGH SCHOOLS, CHICAGO TWENTIETH CENTURY TEXT-BOOKS THE ELEMENTARY PRINCIPLES OF CHEMISTRY BY A. V. E. YOUNG PROFESSOR OF CHEMISTRY IN NORTHWESTERN UNIVERSITY "Not blind To worlds unthought of till the searching mind Of Science laid them open to mankind " WORDSWORTH NEW YORK D. APPLETON AND COMPANY 1901 COPYRIGHT, 1900 BY D. APPLETON AND COMPANY PREFACE THIS book on elementary chemistry is based on the plan which, without essential modification, the writer has been following for thirteen years with his classes beginning the subject. Its inception was under the stimulating sugges- tion of the late Professor Josiah P. Cooke, of Harvard Uni- versity, an early advocate of what has been called the quan- titative method in teaching the subject even in elementary courses. The temerity of the writer in thus offering to his fellow-students and teachers the outcome of his personal experience is due to the hope, perhaps egotistic, that he may contribute, in howsoever small a measure, to making practicable and serviceable that which, he enthusiastically believes, is both scientifically and pedagogically an improve- ment on the older and still largely prevailing method. In the preparation of the book two things have been assumed : first, that sufficient laboratory facilities are pro- vided ; and, second, that the teacher has information in the subject beyond that which the book itself supplies. The laboratory work by the student is made a large and essen- tial part of the exposition, and it is generally assumed that he has performed the experiment illustrative of a topic before he gives attention to the fuller presentation of the same in the text. 84369 vi ELEMENTARY PRINCIPLES OF CHEMISTRY The writer fully appreciates the increased demand made by such a course on the teacher's time and strength, and he has therefore prepared to accompany this book a teach- er's aid, in which he has put whatever he has been able of suggestions which may be helpful and labor-saving to the teacher, but which do not suitably find place in a book designed for the use of students. A. V. E. Y. EVANSTON, ILL., June, 1900. FOKEWOKD TO THE STUDENT You are about to enter upon the study of a branch of natural science, perhaps the first to which experience has brought you. If I am to have the privilege, shared with your teacher, of being your guide in this, I beg the addi- tional privilege of addressing a word to you, at the outset of your study, in a purely personal manner and free of the formality of authorship. As a lover of nature, whether in the aspect revealed by science, or in that which lies open to him who has eyes to see and soul to feel, and as one to whom science study has brought much genuine enjoyment, I would that to you also this study might be not a task only, even though interesting and profitable, but a source of real pleasure. A large portion of your time will be given to experi- ments. They may seem tedious at times, and may test your patience ; but they are designed to teach, not to amuse, and, in order to teach, they must be performed studiously and thoughtfully. It may be hoped that you will find a pleasure added to mental activity in the accompanying use of the hands. Again, your experiments are generally to be thought of as giving you not proof of laws, but illustration, as aid to clear ideas. In my own classes, I like that the student vii viii ELEMENTARY PRINCIPLES OF CHEMISTRY should, so far as practicable, get the first notion of a topic through his own observation, instead of being told in pre- vious explanation what he is to see. In this way, I believe, there may come to him just a taste of the keen pleasure of discovery that comes to the investigator and pioneer. I would reckon one opportunity of profit from a course in physical science as lost to that student who fails to acquire some increased measure of respect for the teachings of experience, which should serve him to good purpose, if rightly applied, in the daily conduct of life. Moreover, he should learn respect for material things, finding in them not alone facts, useful of application in practical life, and gratifying to the inborn curiosity of the human mind to know, but likewise things worthy of oft- repeated contemplation, and gratifying to the- intellect to admire. The lover of nature is not content with one view of a beautiful landscape; nor is the lover of art content with one look at the handiwork of a great painter or sculptor, but he returns with fresh joy to contemplate the beauty which it embodies. Yet this is the work of the human hand and intellect. Likewise the student of natural science, I am sure, may, if he will, add much to the pleasure of his work and to its ennobling influence by finding in the mate- rial things revealed by his science the handiwork and em- bodied thought of the mightiest of creators, worthy of repeated contemplation and productive of the noblest pleasure. " God is the Perfect Poet, Who in creation acts his own conceptions." ROBERT BROWNING, in Paracelsus. A. V. E. Y. SCHEME OF TOPICS [The numbers following the topics refer to marginal numbers in the text, which are the same for the same topics in both Part I and Part II.] CHAPTER I. Nos. 1-33 INTRODUCTORY In Physical Phenomena two things distinguished : Matter and Energy. Different kinds of Matter = Substances. Transformation of Substances. Definition and Identification of Substances by Physical and Chemical Properties A. Physical Properties (with special reference to Identification) : 1. Taste. 2. Odor. 3. Relative Hardness. 4. Form (crystalline). 5. Mass (measured by weight). 6. Volume. 7. Specific Gravity, No. 7. (Comparative weight.) 8. Behavior toward Light : As to Color, Luster, Transparency, Opacity. 9. Behavior toward Electricity. 10. Behavior toward Magnetism. a. Solids. Fusion, solidification, 10. b. Liquids. Boiling, 11. 11. Behavior toward Heat Distillation, 12. Sublimation, 12. c. Gases. Burning or combustion. Chemical change. x ELEMENTARY PRINCIPLES OF CHEMISTRY B. Chemical Properties, 13. Transformation of Substances. Behavior toward Other Substances. C. Physico-Chemical Properties. 1. Allotropism, 13. 2. Solution, 21. Study of Chemical Change, 15-18. Additional Physical Properties Crystallization, 21/ 4 . Amorphism, Polymorphism, Isomorphism, 21/ 6 . Water of Crystallization, 21/ 7 . Efflorescence, 21 / 8 . Deliquescence, 21/ 8 . (Hygroscopic.) Heat of Solution, 22. Melting Point, 23. Freezing Point. Surfusion. Boiling Point, 24. Conditions affecting. Chemistry is loth : A. Static. Description, Identification, Classification of Definite Chemical Substances. Qualitative Analysis, 26. Quantitative Analysis, 26. B. Dynamic, 26. Factors Products Chemical Change, or of -{ Reaction, or Laws Agency i. e., VM, . /-.i_- i * no -.L ^ Interaction. Chemism or Chemical Affinity Energetics All Substances classified into : a. Elements, 27, and b. Compounds (Mechanical Mixtures, 28). Compounds classified (in part), 30 : Acids, Bases, Salts. Chemical Changes classified (in part), 31 : a. Analytic (Decomposition). &. Synthetic (Combination). c. Metathetic or Mixed (Exchange). Substitution. ' I SCHEME OP TOPICS xi CHAPTER II. Nos. 34-60. THE FUNDAMENTAL QUANTITATIVE LAWS OF CHEMICAL ACTION 1. The Law of Persistence of Mass (Lavoisier), 34. 2. The Law of Fixed Proportions (Proust), 37. 3. The Law of Multiple Proportions (Dalton), 40. 4. The Law of Equivalent Proportions (Richter and Wenzel), 41. Corollary I. Equivalent and Combining Weights of the Elements, 42. Corollary II. Combining Weights of Compounds, 45. Corollary III. Active Masses, the Multiples of Combining Weights, 46. 5. The Law of Gas-volumetric Proportions (Gay-Lussac), 47. Some Illustrations of Relative Quantities in Combination, 49. 6. The Law of Persistence of Energy applied to Chemical Phenomena, 50. The Law of Constant Heat Summation (Hess), 55. CHAPTER III. Nos. 61-65 1. The System of Combining Weights, 61. 2. The System of Notation, 62. , (a) For Elements. (b) For Compounds. 3. Chemical Equations, 63. 4. Stoichiometry, 64. 5. Chemical Nomenclature, 65. Problems. CHAPTER IV. Nos. 66, 67 RELATION BETWEEN VOLUME, PRESSURE, AND TEMPERATURE OF GASES I. The Law of Boyle. Relation between Volume and Pressure of Gases. Corollary. II. The Law of Charles. Relation between Volume and Temperature of Gases. Corollary. Note I and II. xii ELEMENTARY PRINCIPLES OP CHEMISTRY CHAPTER V. Nos. 68-141 RELATION BETWEEN EQUIVALENT AND COMBINING WEIGHTS AND CER- TAIN SPECIFIC PROPERTIES 1. The Law of Gay-Lussac, 71-90. Relation between Equivalent and Combining Weights of Gases, Elementary and Compound, and their Specific Gravities. Note I, 91. Note II, 92. 2. The Law of Dulong and Petit, 95-102. Relation between Equivalent and Combining Weights of Ele- mentary Solids and their Specific Heats. Corollary, 103. 3. The Law of Mitscherlich, 107. Relation between Composition, and hence Combining Weight, and Specific i. e. Crystalline Form. 4. The Law of Raoult (1), 109-123. Relation between Combining Weights of Solutes and Specific De- pressions of the Freezing Point in Specified Solvent. 5. The Law of Raoult (2), 124-135. Relation between Combining Weights of Solutes and Elevations of Boiling Temperature in Specified Solvent. Summary, 136. CHAPTER VI. Nos. 142-156 METHOD OF DETERMINING EQUIVALENT (144) AND COMBINING WEIGHTS (150) OF ELEMENTS AND FORMULAS (153) OF COMPOUNDS Problems, 156. CHAPTER VII. Nos. 157-186 THE ATOMIC THEORY : ITS FUNDAMENTAL ASSUMPTIONS .1. The Molecular Constitution of Matter, 163. 2. The Kinetic Theory of Gases, 164. 3. Avogadro's Hypothesis, 165. 4. The Chemical Definition of a Molecule, 166. 5. As to Molecular Weights, 167. 6. The Divisibility of the Molecule of a Compound, 168. 7. The Divisibility of the Molecule of an Element, 169. 8. The Atom defined, 171. SCHEME OF TOPICS xiii 9. As to Atomic Weights, 173. 10. As to Heat Capacity of Atoms, 175. 11. As to Raoult's Laws, 176. 12. As to Structure of Molecules, 177. Isomers (179), Polymers (179), Metamers (180). 13. As to Space Relations of Atoms, or Stereo-Isomerism, 183. CHAPTER VIII. Nos. 187-641 RELATION BETWEEN THE PROPERTIES OF THE ELEMENTS IN GENERAL AND THEIR COMBINING WEIGHTS Description of the First Twenty-Jive .Elements and some of their Compounds Of the Elements Collectively, 188. . Hydrogen, 200. \^>. Nitrogen, 306. 2. Lithium, 210. Oxides, 312-325. 3. Glucinum, 216. Nitric Acid, 326. 4. Boron, 220. Nitrates, 333. Borax bead, 228. Ammonia, 334. ^ 5. Carbon, 234. Other Compounds, 339. 5a. Carbon Dioxide, 259. Relation to Living Things, 5b. Carbonates, 269. 344. 5c. Carbon Monoxide, 272. /^7. Oxygen, 350. 5d. Hydrocarbons, 274. Ozone, 358. 5e. Flame, 284, Part II. Hydrogen Dioxide, 366. 5f. Petroleum, 293. Water, 368. 5g. Coal, 298. Natural Waters, 371. 5h. Coal Gas, 301. Purification of Water, 378. 5i. Destructive Distillation of Wood, 303. / B. The Atmosphere, 388. C. Argon, 396. . Other New Elements in the Atmosphere, 398. V 8. Fluorine, 401. 9. Sodium, 407. 9a. Sodium Chloride, 411. 9b. Sodium Carbonate, 413. 9c. Sodium Hydroxide, 419. D. General Survey, 422. E. Valence, 434. Review Problems, 441/|. 1 ELEMENTARY PRINCIPLES OF CHEMISTRY 15. Chlorine, 527. 15a. Oxides and Acids of, 536. 15b. Bromine and Iodine, 541A. 15c. Manufacture of Chlo- rine and Bleaching Powder, 542. 16. Potassium, 549. 10. Magnesium, 442. 11. Aluminium, 448. lla. Manufacture of, 462. 12. Silicon, 466. ' 12a. Uses of Silicates, 476. 13. Phosphorus, 484. 13a. Oxides and Acids of, 492. 13b. Other Compounds of, 496. 13c. Manufacture of Phos- phorus and Matches 498. ] 14. Sulphur, 507, 14a. Compounds with Hy- drogen and Chlo- rine, 510. / 14b. Oxides and Acids of, 514. 14c. Manufacture of Sul- phuric Acid, 523. F. Gunpowder and Some Other Explosives, 555. 17. Calcium, 577. 17a. Mortar and Other Cements, 585. Review Problems, 588/i. G. General Survey, 589. 21-25. Chromium, Manganese, Iron, Nickel, Cobalt, 591. 23a. Commercial Iron, 607. H. The Law of Periodicity, 620. LIST OF ELEMENTS IN ALPHABETICAL ORDER, 643. LIST OF ELEMENTS IN NATURAL ORDER, 644. THE ELEMENTARY PRINCIPLES OF CHEMISTRY PAR T I CHAPTER I INTRODUCTION Matter, energy, substances. In physical phenomena the 1* student of nature has learned to distinguish two things, Matter and Energy. Matter is all that which occupies space, has weight, and is acted upon by various agents ; for example, earth, water, air, iron, wood, and salt. Energy is that which brings about change in things ; change in position that is, motion as when a body falls to the earth; or change in nature, as when iron rusts, or wood burns, or water freezes. Energy is therefore the general name, not for any particular agent, but for all agents that bring about the specified changes. Thus, gravitation, heat, electricity, magnetism, and light are only different forms or conditions of energy. Likewise, the very many things familiarly recognized, such as water, air, coal, copper, silver, gold, and others not so familiar, are simply different kinds of matter ; to designate these is used the term, substances. Substances are not permanent. Now, it is readily recog- 2 nized as a fact of common observation, that substances, in many instances at least, are not permanent, but are subject * The marginal numbers are the same for the same topics in both . 2 ELEMENTARY PRINCIPLES OF CHEMISTRY to change ; and furthermore, that the changes often result in the disappearance of the original substances. Thus water may lose its liquid form and become solid, in which condition we call it ice ; or it may become gaseous, in which condition we call it steam or vapor. Wood rots and iron rusts, and in so doing both lose many of the properties which make them useful. Wood and coal burn, and thus are converted into something which is neither wood nor coal. A lump of sugar, or of salt, when dropped into water, completely disappears so far as we can see. And many similar observations might be cited. But, evidently, we can not make even the simple observation that iron ceases to be iron when it rusts, or that wood ceases to be wood when it burns, or, in general, that any substance, A, ceases to be A, or becomes some other substance, B, unless we have some means of identifying the substances, A and B, and distinguishing between them ; and these constitute, in part at least, the subject-matter of Chemistry. For present pur- poses, then, we may define Chemistry as that branch of physical science which deals primarily with the properties and transformations of substances. Physics, on the other hand, deals primarily with the several kinds of energy and their transformations. The two sciences are therefore closely related, and many phenomena may be considered as properly in one as in the other ; and some must, indeed, be considered in both. Identification of substances. We proceed, therefore, to study somewhat in detail the identification of substances. Broadly speaking, a substance can be identified only by its totality of properties, but in practice it is possible to use selected properties, since we have learned by experience that these are accompanied by the others which make up the total. Of the properties of a substance, we designate as physical those which involve no change in the identity of the substance ; as chemical, those which do involve such change. But here, as often when we attempt to classify INTRODUCTION 3 or define the things of nature, we find those which can not be placed satisfactorily in one or the other class, since they partake of the character of both ; and so we shall find it convenient to recognize properties of a third class, and to designate them as the physico-chemical. 1. Physical Properties Among the physical properties useful for identification 4 may be cited : taste, odor, relative hardness, form (crystal- line), specific gravity (i. e., the relation between weight and volume) ; also behavior toward (1) light, as to color, luster, opacity, etc. ; (2) toward electricity ; (3) toward magnet- ism ; (4) toward heat. The chemical properties may be described, with but few 5 exceptions, as the behavior toward other substances ; while the phenomena of solution and of allotropism may serve as illustrations of physico-chemical properties. Before presenting formal statements and definitions con- cerning these items, it is best that they be studied by prac- tical illustrations ; and so sulphur is taken up first, as an object-lesson in description and identification of substances (see Chapter I, Part II). Following the observations and illustrative experiments made by the student in the labora- tory, the matter now to be presented may be considered as a review and summary of the corresponding topics, with some added information, impracticable of illustration. Taste. Taste is of limited applicability for purposes of 6 identification, and in the chemical laboratory should be used with only the utmost caution, as many substances are extremely poisonous. Odor. Odor is more often serviceable, but likewise must be used with caution. Form. By form in this connection is meant only crys- talline form that is, definable, geometric form, which, as will be seen later, is often characteristic. 2 4 ELEMENTARY PRINCIPLES OF CHEMISTRY 7 Specific gravity. The specific gravity of a substance is the ratio of the weight of some sample of it divided by the weight of an equal volume of some standard substance. Water is taken as the standard for solids and liquids, or, more strictly speaking, water at its maximum density (temperature 4 C.). For gases, both hydrogen and air are used ; the former is preferred in this study, and may be understood as the standard for gases unless it is otherwise specified. Inasmuch as the specific gravity of air referred to hydrogen is 14.40, -it is easy to~pa-sHfrom one scale to the other by the use of this factor. 8 Your experiment with sulphur has given illustration of & method for determining the specific gravity of solids which are not soluble in water. If the solid dissolves in water, some other liquid must be used in which the sub- stance is insoluble, and of which the specific gravity is known. Another method is based on the principle that a body when suspended and immersed in a liquid weighs less than when weighed in air simply, and less by a quantity equal to the weight of the liquid displaced. Since liquids and gases may completely fill their con- taining vessel, their specific gravity may be determined by weighing successively the vessel when empty, when filled with the standard, and when filled with the substance in question. 9 Electrification, The electrification of sulphur by fric- tion, so that it attracts particles of matter, is a minor item. Of the same order also, is the familiar observation that iron is attracted by the magnet. The facts as to color, luster, opacity, or transparency usually appear simply on inspection. 10 Behavior toward heat. The behavior toward heat is of prime importance. Three conditions of matter are recog- nized the solid, the liquid, and the gaseous. These are dependent on the temperature, or on the temperature and the pressure combined. INTRODUCTION 5 Fusion or melting is changing a solid into a liquid by the agency of heat. Freezing, solidification, or congelation is changing into the solid condition from the liquid, some- times from the gaseous, by the withdrawal of heat. The behavior of sulphur, which you have noted, is peculiar in that the continued application of heat after fusion causes the liquid first to become semi-solid and then to regain its fluidity. The temperature at which fusion and solidifica- tion take place is definite and constant for a given sub- stance, and hence offers a valuable feature for identifica- tion. Its determination will be considered later. Boiling or ebullition is the formation in a liquid of bub- 11 bles of its own vapor. Evaporation that is, the conver- sion of a liquid into vapor at its surface only takes place at indefinite temperatures, even much below that of boil- ing. But bubbles of vapor form only in definite condi- tions, and therefore the temperature at which boiling takes place is also a valued feature, which later will be studied in detail. Distillation (for experimental illustration see not only 12 Exps. 11 and 12, but also 20/j and 24/ 5 , and Appendix, 18) is the conversion of a solid or of a liquid into vapor by heat- ing, and this vapor in turn into a solid or liquid by cooling. It is serviceable in separating volatile from non-volatile sub- stances, or substances differing in degree of volatility, and is therefore often a valuable means of purification. That which is separated by distillation and condensed is called the distillate. Some substances may be converted from solid to vapor and back to solid without passing through the liquid condition. This operation is called sublimation, and the condensed portion is called the sublimate. In nature the formation of rain may be considered as distillation on an enormous scale, since water passes from the surface of the sea and other bodies of water into the atmosphere as a vapor, is subse- quently condensed into liquid form, and falls to the earth again as drops of rain. Or, the water vapor of the atmosphere may condense 6 ELEMENTARY PRINCIPLES OF CHEMISTRY directly to the solid snowflake or to the frost upon the window pane, and these we may think of as natural sublimates. Distillation and sublimation are used also in many industrial opera- tions for example, in the manufacture of alcohol, the refining of petroleum, of sulphur, and of iodine. 2. Chemical Properties 13 Essential feature. Returning to the observations with sulphur : you have seen that, if the heating is continued beyond the boiling point, the sulphur finally takes fire, burning with a pale bluish flame ; that in consequence of this change, called combustion, the sulphur disappears ; that a substance appears which is gaseous at the ordinary tem- perature, has a peculiar, stifling odor, changes blue litmus paper to red, and is thus a substance distinctly different from the sulphur itself. This change is a chemical one, its essential feature being the production of a substance other than the original it is a change of identity ; whereas the other changes already noted, electrification, fusion, solid- ification, boiling, sublimation, and distillation, all physical, have left the substance still sulphur. The peculiar product obtained by turning the liquid sulphur at about its boiling point into cold water certainly differs in some of its prop- erties from the original substance. On inspection you hardly recognize it as sulphur, yet on standing^t passes back into the brittle condition without a change in weight, and en burning it yields the same gaseous product that the brittle sulphur does. Such a change, considerable yet not suf- ficient to constitute a change in identity, is called allo- tropic ; and the plastic or amorphous substance and the brittle or crystalline are said to. be allotropic forms of sul- phur. Further illustration of chemical change you have seen in the production of iron sulphide from iron and sulphur, zinc sulphide from zinc and sulphur, hydrogen sulphide from hydrochloric acid and iron and zinc sul- phides. INTRODUCTION 7 Secondary features. You may also note as secondary differences between physical and chemical changes, that the former may involve but a single substance, whereas chem- ical changes must involve at least two, and generally in- volve more than two. There are a few, but only a few, changes which fall under the general form : substance A becomes substance B, or, more briefly expressed, A = B. Therefore the chemical properties of a substance may be generally described as its behavior toward other substances. Again, the physical changes are often reversible by revers- ing the conditions ; thus sulphur changes from solid to liquid and then to vapor by rise of temperature, and then a fall of temperature reverses the changes. On the other hand, if the temperature be raised until the sulphur burns that is, until chemical action takes place and the new substance, sulphur dioxide, is formed this change is not reversible ; that is, the sulphur can not be recovered by cooling the gas. 3. Additional Illustrations of Chemical Change The additional illustrations of chemical change have brought to your attention the four types, which may be given general expression in the following concise equation forms, and which for the dwelt upon in connection with the illustrative experiments : (1) ^r^FTT= AB (compositiop). 15 (2) "AB A + B (decomposition). 16 (3) A + BC = B.+ AC (substitution). 17 (4) AB + CD = AC -f BD (double exchange). 18 4. Additional Illustrations of Physical Properties Solution. We proceed now further to consider some of 21 the physical properties. Solution is the conversion of a solid or of a gas into a liquid by the action of a liquid, 8 ELEMENTARY PRINCIPLES OF CHEMISTRY The term is also applied to the mixing of one liquid with another. The substance dissolved is called the solute, the liquid which effects the solution the solvent, and, unfor- tunately for clearness, the mixture of these two that is, the result obtained by dissolving is called the solution. The phenomenon of solution has practically a limitation, qualitative in character that is to say, some substances dissolve in some solvents and not in others, while some are practically insoluble in all solvents. Thus sugar, common salt, alum, and copper sulphate dissolve in water, but not in carbon disulphide ; while sulphur dissolves in the latter (see Exp. 12/ 2 ), but not in water. And yet those sub- stances commonly reckoned as the most insoluble do, in many instances, dissolve in minute quantity. Thus, for in- stance, sand, limestone, and glass would in ordinary expe- rience be thought insoluble in water, yet they do appre- ciably dissolve. 21/1 Solution plays an important part, often on a vast scale, in the pro- cesses of nature, and in these, water is the great solvent. Water pass- ing through the atmosphere and over or through the earth's crust dis- solves many substances in quantities small perhaps in proportion to the quantity of the solvent, yet large in the aggregate, since the quan- tity of water is enormous. These ultimately, at least in large part, reach the sea, which thus becomes a vast reservoir, not only of water, but also of soluble matter from the solid portion of the earth's crust. This is the most stupendous instance of the phenomenon of solution within our observation. In the vital processes of the plant and of the animal, solution is of great importance, serving to bring the constituents of the food into condition suitable for distribution and assimilation in the different parts of the organism. 21/2 Solution is an operation of prime importance also in the industrial ' arts as well as in the laboratory. It is used not only to separate soluble from insoluble substances and substances of different degrees of solu- bility, and consequently as a means of purification, but also to bring about a condition suitable for some other operation or change. The refining of sugar is one of the innumerable instances. 21/3 The property of solubility or insolubility, or the degree of solubility, is also of great practical value in the identifi- INTRODUCTION 9 cation of substances. The quantity of a given substance which a definite quantity, say one hundred grams, of a given solvent can dissolve is limited, and is dependent on temperature, and, in the case of gases, on pressure also ; but for a definite temperature (and pressure) the quantity is definite. The solubility of solids generally increases with increase of temperature, but in some instances it decreases. The solubility of gases diminishes with increase of temper- ature, and it increases with increase of pressure. A solvent, when it has dissolved the maximum of a given substance, is said to be saturated. Unsaturated solutions of solid or non-volatile substances may be concentrated by evaporat- ing the solvent (see Exp. 21/ 4 ), and the solute may be recovered, often unchanged, by making the evaporation complete (see Exps. 17/i b and 18/A). Crystallization. If saturated or nearly saturated solu- 21/4 tions remain undisturbed, so that slow cooling and evapora- tion may take place, the solid in many instances separates in definite geometric forms, as you saw in your own experi- ments with alum and copper sulphate. This is crystalliza- tion. It may accompany solidification, not only from the state of solution, but also of fusion (seen in the sulphur ex- periments, !S r os. 11 and 12, and 12/i), and even from the gaseous condition (seen in the experiment with iodine, No. 20/i , the sublimate of which is beautifully crystalline). In the crystallization of mixtures each individual substance crystallizes by itself, in its own peculiar form, and therefore crystallization is a most important means of separation and purification. Furthermore, the peculiarities of form that is, the shape and disposition of the faces, the dimension of the angles, etc. are definite and constant characteristics of the substance, and therefore of great value in identifi- cation. .Crystallization on the large scale finds abundant illustration in 21/5 nature, since many constituents of the earth's crust exist in crystalline condition, either in large masses of crystalline structure, like granite, 10 ELEMENTARY PRINCIPLES OF CHEMISTRY marble, etc., or as distinct individual crystals like the diamond, ruby, emerald, and many other valued gems. Many manufactured products also appear in crystalline condition, of which sugar and salt may be cited as examples. 21/6 Substances when they do not show this definite geomet- ric or crystalline form are said to be in the amorphous con- dition. Some substances are capable of crystallizing in two or more distinctly different forms. They are said to be polymorphous, or, if they exhibit only two forms, dimor- phous. This property you have seen in the two crystal- line forms of sulphur (see Exps. 12/ 1? and 12/ 2 ). On the other hand, some instances are found of different substances exhibiting the same crystalline form, and such substances are called isomorphous ; examples are calcium, magnesium, iron, and zinc carbonates. 21/7 Many substances, not all, in crystallizing from water solution, contain water as a constituent, although no evi- dence of this is seen by simple inspection of the crystal. This is termed water of crystallization. It can generally be driven off by heating, and the crystalline structure is thus destroyed, as you have seen with copper sulphate. In some instances the crystal first melts in its water of crystal- lization, as in the case of alum. Substances from which the water has thus been taken and substances which contain no water are said to be dehydrated or anhydrous. Again, 21/8 some substances lose their water of crystallization at the ordinary temperature; such are called efflorescent; exam- ples, sodium carbonate and sodium phosphate. On the other hand, some substances absorb water from the atmos- phere and tend to liquefy in consequence ; these are de- scribed as deliquescent ; examples, zinc chloride (see Exp. 17/i b), calcium chloride, and sodium hydroxide. The term hygroscopic also is applied to substances which thus absorb water, but this does not imply liquefaction. Thus, com- mon quicklime that is, calcium oxide takes water abun- dantly, but retains the solid form. INTRODUCTION 11 Heat of solution. Solution, when solute and solvent 22 are of the same temperature before mixing, is usually, not always, accompanied by reduction of temperature. Thus, common salt dissolving in water lowers the temperature, while sodium hydroxide and hydrochloric acid raise it. Melting point. The melting point of a substance is the 23 temperature at which it melts, and the freezing point is the temperature at which it solidifies ; and, being definite and constant characteristics of a specified substance, they are, as already suggested, useful in description. One would ex- pect the two temperatures to be really identical, yet obser- vation shows that the liquid form is sometimes retained below the melting point. This phenomenon, called sur- fusion, may cause irregularity in the observation of the freezing point. Your experiment gives illustration of one method applicable in such observations. "With a larger quantity of the substance, the bulb of the thermometer may be thrust into the liquid, and the latter be stirred during the solidification. The presence of substances in solution tends to lower the 23/1 freezing point of the solution as compared with that of the pure solvent. This phenomenon will be studied quantita- tively in Chapter V. Boiling point, and circumstances affecting it. The boiling 24 point of a substance, defined in the strictest sense for pur- pose of description, is the maximum temperature which its vapor attains when it is in contact with the boiling liquid and free to escape into the atmosphere ; or, in other words, when it has a pressure or tension equal to that of the sur- rounding atmosphere. If this pressure is increased by con- 24/1 fining the vapor or by rise in the atmospheric pressure, then the temperature of the vapor and of the boiling liquid will be higher ; if the pressure is decreased by removal of the vapor or by fall in the atmospheric pressure, the tem- perature will be lower. For complete definition, therefore, the pressure of the vapor, as well as the temperature, should 12 ELEMENTARY PRINCIPLES OF CHEMISTRY be specified, and the normal atmospheric pressure at the sea level that is, 760 millimeters or 29.92 inches is chosen as the standard pressure. The ordinary variations of the atmospheric pressure, however, have so small effect on the boiling point that the variation may for many pur- poses be ignored. A change of 27 millimeters (or 1.06 inches) in barometric pressure, or of about 1,000 feet in elevation, makes a difference of 1 C. in the boiling point. Thus water, boiling at 100 C. at sea level, boils at 85 C. on the top of Mont Blanc. The facts that increase of pressure raises this temperature, and that decrease of pressure lowers the same, find many practical appli- cations. Thus in extracting gelatine from bones, the solvent power of the water is increased by boiling under pressure ; and in the refining of sugar the concentration of the sirup is greatly facilitated by evap- oration under diminished pressure. The boiling point has been defined, in accordance with scientific usage, as the temperature of the vapor. Yet the term is also used, even in scientific literature, to designate the temperature of the boiling liquid. This is unfortunate, since it is important to discriminate between the two, for the temperature of the liquid is often higher than that of the vapor. Various circumstances cause this. Whatever 24/2 tends to hinder the formation and escape of the bubbles of vapor tends to raise the temperature of the liquid above that of the vapor ; conversely, whatever tends to facilitate the same, tends to make the temperature of the liquid the same as that of the vapor. Thus, water boils more readily in an iron vessel with rough surface than in a glass one with smooth surface. Indeed, in a glass vessel, whose inner surface has been very thoroughly cleaned, water may be raised to a temperature considerably above 100 C. In this condition a grain of sand or a fragment of glass dropped in, or a stirring rod inserted, may cause sudden and violent boiling. This phenomenon of irregular boiling, commonly called in the laboratory bumping, is often very annoying, INTRODUCTION 13 and various expedients are used to prevent it, such as con- stant stirring, or fragments of broken glass, or pumice, or platinum foil. Thus also the presence of dissolved air or other gas tends to facilitate boiling. On the other hand, 24/4 the presence of non-volatile substances in solution tends to raise the temperature of the boiling solution as compared with that of the pu?e solvent ; this fact has important quan- titative relations, which will be considered in Chapter V. 5. Definitions and Classifications Some branches of chemistry. It is suitable to introduce at this point some additional statements and definitions which do not call for special experimental illustration, although it is not to tie expected that the full significance of these statements will be understood at this stage of the course. From what has been already presented, it may be seen that the subject-matter of our study has a twofold aspect. We 26 may give attention primarily to the description and iden- tification and classification of definite chemical substances, and this may be called static chemistry ; or, on the other hand, we may give chief consideration to the changes of sub- stances as changes, and this is the aim of dynamic chemistry. Of the first division are the subjects qualitative analysis and quantitative analysis. The former consists of syste- matic methods for the identification and recognition of sub- stances ; the latter, of methods for determining the quan- tities of substances. Analysis is ultimate when it is con- cerned with the qualitative or quantitative determination of elementary constituents ; and proximate when it is con- cerned with constituents other than elementary. In dynamic chemistry we have to consider such subjects as the factors, products, laws, agency, energetics, etc., of chemical change. For this last phrase the more technical one, chemical reaction or interaction, is commonly used. The factors are those substances in presence before the 14 ELEMENTARY PRINCIPLES OP CHEMISTRY change takes place ; the products, those present after the change. The factors and the products considered collect- ively may be conveniently designated as the system under- going change. 27 Elements and compounds defined. All substances may be classified as elements or as compounds. An elementary sub- stance is one from a given weight of which no other sub- stance has been obtained less in weight than the original. Examples : sulphur, iron, zinc, iodine. A compound sub- stance is one from a given weight of which other substances have been obtained, each less in weight than the original. Examples : sulphur dioxide, iron sulphide (Exp. 13/ 2 a), zinc sulphide, lead iodide (Exp. 15/j), magnesium, zinc, and lead oxides (Exps. 15/ 2 , 15/ 3 , and 15/ 4 ), lead nitrate (Exp. 16/!), zinc nitrate (Exp. 16/ 2 ). 28 It is important to distinguish between a chemical com- pound and a mechanical mixture. The first is homogeneous, and has properties of its own, distinct from those of its constituents ; the second is not homogeneous, and shows only the properties of its components. Thus the powdered zinc and sulphur in your experiment (Exp. 13/j), before the application of heat, is a mechanical mixture, showing only the properties of the zinc and of the sulphur. A micro- scope distinguishes the different color of the zinc and sul- phur particles. Water, dissolving neither, separates them more or less completely by difference in specific gravity. Carbon disulphide dissolves the sulphur and leaves the zinc, while hydrochloric acid dissolves the zinc and leaves the sulphur. But the product of the chemical change that is, the zinc sulphide has its own properties, distinct from those of zinc and those of sulphur. 29 It is customary to divide the elements into the metals and those which are not metals, that is, the non-metals, *and to consider the two classes separately. This classifica- tion has largely lost its original significance and impor- tance, however, and it is preferred, in this presentation, to INTRODUCTION 15 make the classification a secondary matter, and to use the term metal rather in the ordinary non-technical sense, im- plying those properties associated in familiar observation with substances like iron, copper, gold, and silver. But special importance is attached to the question concerning each element whether its combination with oxygen pro- duces a base-forming or an acid-forming substance. The elementary substances in alphabetical list are given in Table X, No. 643. They number seventy-four, and to this com- paratively small number of constituent substances all known forms of matter are reducible. Compounds classified. Three groups of compounds may 30 be defined now ; they are important, although they do not by any means include all compounds. These are the adds, the bases, and the salts. The adds are substances contain- 30/1 ing hydrogen as a constituent, which hydrogen is replace- able by a metal, the product being a salt. When soluble in water, as most of them are, they are sour, change blue litmus to red, combine with bases to form salts, and in con- centrated form are often very corrosive. Examples : hydro- chloric, nitric, and sulphuric acids. A base is a substance 30/2 containing oxygen, often hydrogen also, and always a con- stituent other than these two which in the most common bases is a metal. When soluble in water, they are generally slimy to the touch, bitter, corrosive, change red litmus to blue, and combine with acids to form salts. Examples: ammonium hydroxide, sodium hydroxide, lime, that is, cal- cium oxide, and magnesium, zinc, and lead oxides. The salts are substances formed either by the replacement of 30/3 the hydrogen of an acid by a metal, or by the combination of an acid and a base. The peculiar properties of the acid and the base disappear in the salt, or are neutralized, some- times completely, sometimes only in part. When they are neutralized exactly, so that the salt has no action on lit- 30/4 mus, it is said to be neutral. Examples : ammonium chlo- ride, copper sulphate, zinc sulphate, lead nitrate. When 16 ELEMENTARY PRINCIPLES OF CHEMISTRY all the replaceable hydrogen of the acid is completely and 30/5 exactly replaced by its equivalent of metal, or basic con- stituent, the salt is said to be normal. A normal salt may have neutral, acid, or basic reaction on litmus ; thus, ammo- nium chloride is normal and neutral, alum is normal and has acid reaction, sodium carbonate is normal and has basic reaction. When a salt contains more than the normal 30/6 equivalent of acid, it is said to be an add salt ; such are sodium acid carbonate, and acid sulphate. When it con- tains more than the normal equivalent of base, it is said to be a basic salt ; such is the basic lead acetate. 31 Reactions classified. Chemical reactions may be classi- fied, in part only, as analytic, synthetic, and metathetic, or mixed : analytic, when they change compounds into their constituents ; synthetic, when they change constituents into compounds; and metathetic, when they involve both analy- sis and synthesis, or the exchange of constituents. Thus the conversion of iron and sulphur into iron sulphide is purely synthetic ; of iron sulphide into iron and sulphur is analytic ; while the action of hydrochloric acid on iron sul- phide is metathetic, since the hydrogen and the chlorine of the first and the iron and the sulphur of the second are separated, and then the hydrogen and the sulphur combine, also the iron and the chlorine. When a constituent of a compound is caused to leave it, and another substance appears as constituent in its stead, the process is called substitution (see Exps. 17, etc.). 32 Conclusion. The leading idea of this chapter is the fact . of chemical change that is, the transformation of sub- stances, involving the disappearance of some, and, depend- ent upon this, the appearance of others. This has been called change in identity. Along with this have been pre- sented some definitions and classifications. And in this connection it is interesting and instructive again to note that the application of these to natural phenontena is ofen unsuccessful. Thus it seems a simple matter to define INTRODUCTION 17 chemical changes as has been done ; nevertheless changes are not infrequently encountered, of which it is impossible to decide whether they are chemical or not, because it is impossible to determine whether or not there is a distinctly new substance produced. So also with regard to acids and bases : some substances act in one compound as acid and in another as base, and in some instances there seems as good reason to call it the one as the other. None the less, we may make profitable use of such definitions and classifi- cations ; and at the same time, we do well to bear in mind the fact that nature is not limited by the boundaries of man's thought. CHAPTER II THE FUNDAMENTAL QUANTITATIVE LAWS OF CHEMICAL CHANGE 1, The Law of Persistence of Mass * 34 HAYING now acquired a notion of the method of iden- tifying and differentiating substances, and also of the re- markable transformations which a system of substances may undergo by which other substances are produced, differing entirely from the original in properties, we pass to a closer study of the peculiar quantitative character- istics of these chemical changes. And the most funda- mental fact to be learned is, that by these changes the quantity of matter that is, the mass of the system is neither increased nor diminished. That this is true, does not by any means appear upon the face of things. Numer- ous experiences of everyday occurrence can be cited which would seem to prove that the reverse is true. Thus in the burning of wood, or of coal, or of a candle, one would con- clude from appearances, as people did for many years, that matter is destroyed, or its quantity diminished. If we drop a piece of marble into hydrochloric acid, the marble slowly disappears. On the other hand, in seeing mercury sulpho- cyanate burn, one might suppose that the quantity of mat- ter is largely increased. If, howeVer, it is burned on the 34/1 balance, we see at once that the first supposition is wrong ; that, although it has increased much in volume, the residue which is left on the balance after burning actually weighs less than the original material ; and, if we burn a weighed 34/2 taper and collect and weigh the products of combustion, 18 QUANTITATIVE LAWS OF CHEMICAL CHANGE 19 we find that the products weigh more than the taper con- sumed. Likewise, if we weigh the lump of marble and the 34/5 acid before bringing them together and afterward, we learn that loss of weight accompanies the disappearance of the marble. What is the explanation of these apparent contradictions ? To answer this question we must investigate still further beyond appearances. We note that the disappearance of the marble is also accompanied by effervescence. This im- plies the liberation of a gas, the escape of which must cause loss in weight. We operate so as to retain the gas, and we r find neither loss nor gain in weight. Likewise we may learn that when the mercury sulphocyanate and the taper burn, there is combination of the oxygen of the air, an invisible gas, with the material of the combustible ; and that other gases are produced, also invisible, which by escaping cause loss of weight ; and, furthermore, that these gaseous products, in the case of the taper, weigh more than the taper consumed by just the weight of the oxygen which has entered into combination. The error of the earlier conclusions, therefore, has come from overlooking some of the substances involved in the change. It is only by taking them all into account that the truth has been reached. This great law is known as The Indestructibility of Mat- 34/6 ter, or, better, as The Persistence * or Conservation of Mass. It may be stated as follows : In any system of substances undergoing chemical change the mass of the entire system remains constant, or the total mass of the factors is equal to the total mass of the products. This law is found to hold true throughout the whole 34/7 range of experience, absolutely without exception ; but that * " Persistence" is preferred to "Conservation," following the sug- gestion of Herbert Spencer in " First Principles." 20 ELEMENTARY PRINCIPLES OF CHEMISTRY it is based solely on experimental observation, and is not by any means axiomatic, should not be forgotten. 34/8 It was not discovered until the balance was used to give its evidence. And this was done in the main by Lavoisier, a French chemist, in a series of investigations, extending from 1770 to 1780, which showed the error of previous con- ceptions, and which are regarded as the beginning of modern chemistry. 34/9 This law has its parallel in that other great generaliza- tion of more recent discovery, which lies at the founda- tion of physical science, viz. : The Law of the Conserva- tion or Persistence of Energy, which is that in any system^ of bodies, mutually interacting, the total energy of the system remains constant. 2. The Law of Fixed or Definite Proportions 37 Analysis of your experiment with ammonium hydroxide and hydrochloric acid (No. 37) shows, within the limit of accuracy attainable in the given conditions, that the masses .of these two substances which combine to form the salt, ammonium chloride, bear a ratio to each other which is constant in the three parts of the experiment ; that an ex- cess of either constituent above this ratio has no effect upon the quantity of salt ^btained nor upon its properties. In other words, the three portions of salt obtained are dif- ferent samples of the same definite substance and contain the same relative, quantities of the constituents, ammonium hydroxide and hydrochloric acid. 37/4 The fact here illustrated is found to be a general one, and to bear the test of the most refined experimental methods without established exception. It is formulated in The Law of Fixed Proportions, which may be thus stated : 3 7/5 Any sample of a definite chemical substance, not elementary, is composed of the same constituents, combined in the same relative quantities, as any other sample of the same substance. ANTOINE LAUBENT LAVOISIEK B. Paris, 1743. D. on the scaffold, 1794. (See Nos. 34/ 8 , 235, 259, 306, 388.) QUANTITATIVE LAWS OF CHEMICAL CHANGE 21 The conception of the fixity of ratios was in the minds 37/6 of chemists more than one hundred years ago. Bergmann was apparently guided by it, and Lavoisier distinctly for- mulated it, but its truth was not admitted without dispute. It was called in question in 1799 by Berthollet, a famous French chemist. This led to a long discussion, one of the most remarkable in the history of the science, between him , and another French chemist, named Proust, who undertook the defense of the law. The ideas of the latter prevailed, and, by 1806, the truth of the proposition was generally admitted ; although even in late years there have appeared suggestions that the law may be subject to variations within very narrow limits. In view of this history and of the nature of the law itself, the same degree of absoluteness should not be claimed for it as for the first law. It is to be noted that the Law of Fixed Proportions 37/7 affirms the intimate relation between properties and com- position. With identity of properties there is identity of composition. It does not, however, follow from this, nor is it indeed true, that with identity of composition there is identity of properties. There are many substances not identical, which are nevertheless composed of the same constituents combined in the same quantitative ratio. To the Law of Fixed Proportions may be given an 37/8 Alternative Statement, somewhat broader and more suggest- ive of the dynamic conception, thus : In any system of substances undergoing chemical change, the active masses that is, the quantities actually talcing part in the change, both of factors and of products hear a fixed ratio to each other, always the same for a specified change. 3. The Law of Multiple Proportions The third characteristic of chemical change is perhaps 4-0 even more striking than the two already studied, and is of very great significance. It was noted by Proust, and still 22 ELEMENTARY PRINCIPLES OF CHEMISTRY earlier by Richter, that some substances combine in more than one ratio, producing distinctly different substances. They failed, however, to note the multiple relation of the varying quantities. This was due, perhaps, in part to the inaccuracy of their quantitative determinations. It was left to Dalton, a teacher and chemist living in Manchester, , England, to discover this important relation and to an- nounce it about the year 1804. The law is illustrated in your experiments with mercury and iodine (Exps. 40, etc.), wherein it is seen that these substances combine in two different ratios, producing two distinctly different substances, and that for the same weight of mercury the weight of iodine in one is just twice that in the other. The general fact is known as The Law of Multiple Pro- portions, and may be stated thus : If two substances combine in more than 'one ratio, form- ing distinctly different products, and if in such cases the quantity of one constituent is reckoned as constant, then the varying quantities of the other constituent are in the ratio of small whole numbers. 40/6 It was the discovery of this law, based, it is true, on very few observed instances, that led Dalton to the inven- tion of the atomic theory in its modern form a theory which in importance is second to none in the science of chemistry. 40/7 NOTE. The law holds without exception, although there are many instances in the so-called organic substances, com- pounds of carbon with other elements, in which the rela- tion seems less simple than is stated in the law, but they do not constitute valid exceptions to the law for reasons which would hardly be understood at this stage of the course. 40/8 The multiple relation holds for quantities measured by volume as well as by weight, if the substances are in the gaseous condition. QUANTITATIVE LAWS OF CHEMICAL CHANGE 23 4. The Law of Equivalent Proportions * It has been seen in your experiments, Nos. 41/! and 41/ 2 , 41 that 24 grams of magnesium and 65 grams of zinc combine respectively with 16 grams of oxygen approximately. De- fined with the utmost accuracy, the statement is that 24.1 grams of magnesium and 64.91 grams of zinc combine re- spectively with 15.88 grams of oxygen. It is also a fact, although not shown in your experiments, that these same masses of magnesium and zinc combine respectively with equal quantities of chlorine, namely, (2 X 35.18) grams, and with equal quantities of iodine, namely, (2 X 125.89) grams. Then in your experiments, Nos. 41/ 3 and 41/ 4 , it is seen that these masses likewise displace equal quantities of hy- drogen from an acid. This value, most accurately deter- mined, is 2 grams. Now, magnesium and zinc do not com- bine with each other, but 2 grams of hydrogen exactly combine with 15.88 grams of oxygen in the formation of water ; and one gram of hydrogen combines exactly with 35.18 grams of chlorine, and with 125.89 grams of iodine. Still another phase of this phenomenon is seen in the 41/6 following facts : 31.83 grams of sulphur combine with 64.91 grams of zinc ; also, 15.88 grams of oxygen with 64.91 grams of zinc ; and sulphur and oxygen combine with each other in the ratio of 31.83 : (2X15.88) in sulphur dioxide and 31.83 : (3x15.88) in sulphur trioxide. Generalization upon the many facts of this nature gives * Inasmuch as the experimental study of this law involves the measurement of gas-volume, the instructor may prefer to introduce * the laws of Boyle and of Charles, Chapter IV, at this point. The writer prefers to give them simply as arbitrary rules and to study them later, rather than to interrupt the logical development of this chapter. 24 ELEMENTARY PRINCIPLES OF CHEMISTRY 41/7 The Law of Equivalent Proportions, to which there seems to be no exception ; it may be stated thus : The masses of two or more substances, A,B,C, etc., which combine respectively with a constant mass of another sub- stance, M, are also the masses or bear the relation of small multiples to the masses of A, B, C, etc., which may combine respectively with a constant mass of any other substance, X, or which may be chemically substituted respectively for a constant mass of a substance, X, or which may combine with each other. We may conceive that the masses of two substances, A and B, produce equal chemical effect when they combine respectively with a constant mass of the substance, X ; like- wise when they displace a constant mass of the constituent, C, from the compound, C D ; also when they combine with each other to form the compound, A B. By means of this conception the law may be given a more general expression 41/8 in An Alternative Statement, thus : Those active masses of substances which produce equal chemical effect in a given reaction are also the active masses or bear the relation of small multiples to the active masses of the same substances which produce equal effect in any other reaction in which they take part. 41/9 The first notion of the remarkable facts generalized in the fourth law, without by any means a conception of their full extent or of their significance, is credited by some to a German chemist, Eichter, and to the year 1792, and by others to Wenzel, and the earlier date, 1777. Corollaries. The following important corollaries are de- ducible from the fourth law, taken in connection with those which precede : 42 COROLLARY I. It is possible to determine by experiment what mass of every elementary substance combines with one gram of hydrogen (the lightest substance specifically and chemically) or with that quantity of some other substance which, in its turn, combines with one gram of hydrogen QUANTITATIVE LAWS OF (practically 7.94 grams of oxygen, or 85.18 grams of chlo- rine] ; or, what mass displaces one gram of hydrogen from some specified compound. And, moreover, the constituent masses of the elements in any compound, and their active masses in any reaction, stand in the ratio of these quan- tities or of multiples of the same by a whole number. Definition. These fundamental quantities, specified in 4:3 Corollary I, are called the equivalent weights or equivalent masses of the elements. Now, many of the elements com- bine with hydrogen or with oxygen, or with chlorine in more than one ratio. These substances would have, therefore, two or more values which would answer the definition of equivalent weight. But since the relation of multiples holds between these values, as expressed in Law 3, the fact that there may be more than one value does not interfere with the truth of Corollary I. Great significance, both practical and theoretic, attaches 44 to the equivalent weights and to the facts formulated in this corollary. And the whole system of expressing mass relations in chemical phenomena is based upon the use of the equivalent weight, or of a multiple of the same, as a chemical unit, peculiar to each element, and upon the ex- pression of the relative active mass in terms of this unit, with a coefficient which is always a whole number. That such a system is possible, is clearly a consequence of the foregoing laws. That one equivalent rather than another is used, when two or more are possible, or that a multiple instead of the equivalent itself is used as the basal unit, is a matter of choice, determined conventionally by the applica- tion of certain principles or of certain theories. For the reasons which are presented in detail in Chapter V, the equivalent weights of certain of the elements have been multiplied by a small whole number, two, three, four or five, and the values thus obtained together with those of the equivalent weights which are left unmodified, all of them determined with the utmost available accuracy, constitute 26 ELEMENTARY PRINCIPLES OF CHEMISTRY the elementary combining weights, which are made the basal units, just referred to, for expressing the quantitative mass relations in chemical phenomena. The elements with their equivalent weights, the factors, and the combining weights are given in Table XI, Ko. 644, in natural order i. e., in the order of increasing combining weights. 45 COKOLLARY II. It is evident that the combining mass of a compound must be the sum of its elementary constituent combining masses. For example : Using round numbers simply, the combining weight of water is 18 grams, since it contains hydrogen and oxygen in the ratio of 2 grams to 16 grains that is, two combining weights of the former to one combining weight of the latter ; of ammonia is 17 grams, since it contains nitrogen and hydrogen in the ratio of 14 grams to 3 grams that is, one combining weight to three combining weights ; of carbon dioxide is 44 grams, since it contains carbon and oxygen in the ratio of 12 to 32 that is, one combining weight to two combining weights. 46 COKOLLARY III. It is possible, therefore, to represent the relative active masses of all substances, in any reaction what- soever, as multiples by whole numbers of their respective com- bining masses. 5. The Law of Gas- volumetric Proportions 47 The experiments (No. 47, etc.) with nitrogen diox- ide and oxygen, designed to illustrate this law, are only partly satisfactory, since the volume of the product can not be measured in the conditions of the experiment, the substance, nitrogen tetroxide, being soluble in wa- ter. Additional illustrations are found in the following facts : 1 gas-vol. of hydrogen and 1 gas-vol. of chlorine (sum = 2) form 2 gas-vols. of hydrochloric acid. 2 gas-vols. of hydrogen and 1 gas-vol. of oxygen (sum = 3) form 2 gas-vols. of water. QUANTITATIVE LAWS OF CHEMICAL CHANGE 27 3 gas-vols. of hydrogen and 1 gas-vol. of nitrogen (sum = 4) form 2 gas-vols. of ammonia. In these is revealed the surprising but none the less unmistakable fact, based solely on experimental obser- vation, that the gas-volume of a compound is not always equal to the sum of its constituent volumes, but may be less. Facts of this nature are generalized into The Law of Gas-volumetric Proportions as follows : If the constituents are gaseous or volatile their combining volumes, as gases, are in the ratio of small whole numbers, and the sum of these volumes stands in simple relation to the volume of the resulting compound, if the latter is gaseous or volatile. In general, the compound occupies two unit gas- volumes. Note carefully that the law is applicable only to gas- 47/5 volumes ; no such relations hold between combining volumes other than gaseous. The accuracy of the law is of course limited by the ac- 47/6 curacy in the measurement of gas-volumes. There has been much dispute as to the generality of the rule that the com- pound occupies relatively two unit volumes. In several instances its volume is undoubtedly larger, but it has been proved with much ingenuity that these exceptional volumes are due to the decomposition (dissociation) of the com- pounds in the conditions of observation, so that the observed volume is really the volume of a mixture of the constituents, and not of the compound solely. The discovery of the law is due mainly to Gay-Lussac, 47/7 who with Humboldt determined in 1805 the gas-volumetric composition of water, and subsequently extended his inves- tigations to other substances. The large significance of this law in the interpretation of chemical phenomena was not realized, however, until a considerably later date. 28 ELEMENTARY PRINCIPLES OF CHEMISTRY TABLE I 49 As aid to a clearer understanding of these laws, the illustrations which have occurred in the experiments, and some additional ones, are tabulated as follows : 1 gram of hydrogen (H) combines with 35.2 grams of chlorine (Cl), forming 36.2 grams of hydrochloric acid (HC1). 79.3 grams of bromine (Br), forming 80.3 grams of hydrobromic acid (HBr). 125.9 grams of iodine (I), forming 126.9 grams of hydriodic acid (HI). 2 grams of hydrogen combine with 15.9 grams of oxygen (0), forming 17.9 grams of water (H 2 0). 31.8 grams of sulphur (S), forming 33.8 grams of hydrosulphuric acid (H a S). 3 grams of hydrogen combine with 13.9 grams of nitrogen (N), forming 16.9 grams of ammonia (NH 8 ). 4 grams of hydrogen combine with 11.9 grams of carbon (C), forming 15.9 grams of methane (marsh gas) (CH 4 ). 13.9 grams of nitrogen (N) combine with 7.94 grams of oxygen (0), forming 21.8 grams of nitrous oxide (nitrogen monoxide) (N 2 0). 15.9 grams of oxygen (0), forming 29.8 grams of nitric oxide (ni- trogen dioxide) (NO). 23.8 grams of oxygen (0), forming 37.7 grams of nitrogen trioxide (N 2 3 ). 31.8 grams of oxygen (0), forming 45.7 grams of nitrogen tetroxide (N 2 4 ). 39.7 grams of oxygen (0), forming 53.6 grams of nitrogen pen- toxide (N 2 6 6 ). 198.5 grams of mercury (Hg) combine with 125.9 grams of iodine (I), forming 324.4 grams of mercurous iodide ' (Hgl). 251.8 grams of iodine (I), forming 450.3 grams of mercuric iodide (Hgl a ). QUANTITATIVE LAWS OF CHEMICAL CHANGE 29 15.9 grams of oxygen (0) combine with 24.1 grams of magnesium (Mg), forming 40.0 grams of magnesium oxide (MgO). 55.6 grains of iron (Fe), forming 71.5 grams of ferrous oxide (FeO). 63.1 grams of copper (Cu), forming 79.0 grams of copper oxide (CuO). 64.9 grams of zinc (Zn), forming 80.8 grams of zinc oxide (ZnO). 198.5 grams of mercury (Hg), forming 214.4 grams of mercuric oxide (HgO). 205.3 grams of lead (Pb), forming 221.2 grams of lead oxide (PbO). 31.8 grams of sulphur (S) combine with 24.1 grams of magnesium (Mg), forming 55.9 grains of magnesium sulphide (MgS). 55.6 grams of iron (Fe), forming 87.4 grams of ferrous sulphide (FeS). 63.1 grams of copper (Cu), forming 94.9 grams of copper sulphide (CuS). 64.9 grains of zinc (Zn), forming 96.7 grams of zinc sulphide (ZnS). 198.5 grams of mercury (Hg), forming 230.3 grams of mercuric sul- phide (HgS). 205.3 grams of lead (Pb), forming 237.1 grams of lead sulphide (PbS). 31.8 grams of sulphur (S) combine with 15.9 x 2 grams of oxygen, forming 63.6 grams of sulphur dioxide (SO*). 15.9 x 3 grams of oxygen, forming 79.5 grams of sulphur trioxide (SO.). 2 grams of hydrogen are displaced from hydrochloric acid by 24.1 grams of magnesium, forming magnesium chloride (MgCl a ). 55.6 grams of iron, forming ferrous chloride (FeCl 2 ). 63.1 grams of copper, forming cupric chloride (CuCl a ). 64.9 grams of zinc, forming zinc chloride (ZnCl a ). 198.5 grams of mercury, forming mercuric chloride (HgCl a ). 205.3 grams of lead, forming lead chloride (PbCl a ). 36.2 grams of hydrochloric acid (HC1) combine with 16.9 grams of ammonia (NH 3 ), forming 53.1 grams of ammonium chloride (NH 4 C1). 63.6 grams of S0 2 combine with 15.9 grams of 0, forming 79.5 grams of sulphur trioxide (S0 3 ). 30 ELEMENTARY PRINCIPLES OF CHEMISTRY 79.5 grams of S0 3 combine with 17.9 grams of H a O, forming 97.4 grains of sulphuric acid (H 2 S0 4 ). 59.6 grams of NO combine with 31.8 grams of 0, forming 91.4 grams of nitrogen tetroxide (N 2 4 ). 27.8 grams of CO combine with 15.9 grams of 0, forming 43.7 grams of carbon dioxide (C0 a ). 49/1 In the following, the volumes are gaseous : 1 liter of hydrogen (H) combines with 1 liter of chlorine (Cl), forming 2 liters of hydrochloric acid (HC1). 1 liter of bromine (Br), forming 2 liters of hydrobromic acid (HBr). 1 liter of iodine (I), forming 2 liters of hydriodic acid (HI). 2 liters of hydrogen combine with 1 liter of oxygen, forming 2 liters of water (vapor) (H 2 0). 3 liters of hydrogen combine with 1 liter of nitrogen, forming 2 liters of ammonia (NH 8 ). 1 liter of sulphur (vapor) combines with 2 liters of oxygen, forming 2 liters of sulphur dioxide (S0 2 ). 3 liters of oxygen, forming 2 liters of sulphur trioxide (S0 8 ). 2 liters of nitrogen combine with 1 liter of oxygen, forming 2 liters of nitrous oxide (N 2 0). 1 liter of nitrogen combines with 1 liter of oxygen, forming 2 liters of nitric oxide (NO). 2 liters of oxygen, forming 2 liters of nitrogen peroxide (N0 2 ). 2 liters of sulphur dioxide (S0 2 ) combine with 1 liter of oxygen, forming 2 liters of sulphur trioxide (S0 8 ). 4 liters of nitric oxide (NO) combine with 2 liters of oxygen, forming 2 liters of nitrogen tetroxide (N 2 4 ). 2 liters of carbon monoxide (CO) combine with 1 liter of oxygen, forming 2 liters of carbon dioxide (C0 9 ). QUANTITATIVE LAWS OF CHEMICAL CHANGE 31 6. The Law of Persistence, or Conservation of Energy, applied to Chemical Phenomena Heat Disturbance in Chemical Reactions Chemical changes are generally accompanied by changes 50 in energy. You have observed in your experiments several instances in which they have been accompanied by the liberation of heat, sometimes sufficient to raise the mass to incandescence that is, to a red or white heat, as in the com- binations of sulphur with iron and with zinc ; sometimes, less noticeable, as in the action of mercury and iodine ; and sometimes, still less marked, as in the reaction of ammonium hydroxide and hydrochloric acid, and in the action of the latter on magnesium and on zinc. Likewise the reaction between nitric oxide and oxygen liberates heat, although, in the conditions of your experiment, you could not easily observe it. Now, heat is one of the kinds of energy, and the im- mediate agent which causes chemical change is another kind. Energy has been defined as that which brings about changes, or, in the more common wording, it is the power of doing work. The Law of the Persistence of Energy affirms * (compare 34/ 9 , Part I) that, although the various kinds of energy may be transformed from one kind into another, the total energy can not be thereby increased nor diminished. Inasmuch, therefore, as heat, not before evident, appears when iron and sulphur become iron sulphide, and when mercury and iodine become mercuric iodide, it must be that iron sulphide and mercuric iodide as compared with their constituents have lost at least as much energy as is equivalent to the heat produced. The energy which is lost is the chemical energy or chemism, as it is sometimes called. It 51 * This law was announced by J. A. Mayer in 1842, greatly developed by Helmholtz in 1847, and experimentally tested by Joule in 1850. 32 ELEMENTARY PRINCIPLES OF CHEMISTRY may be that the chemical energy which is lost does not appear entirely as heat ; nevertheless, the quantity of heat liberated is constant for a specified change under specified conditions. 52 Again, it follows from this great law that, in order to reverse this change that is, to produce from iron sulphide and mercuric iodide, iron and sulphur, and mercury and iodine the quantity of energy lost in the first change must be restored to the system in one form or another. 53 The quantity of heat liberated in the original reactions is called the heat of formation of iron sulphide and of mer- curic iodide from their constituents respectively. But heat is not always liberated in the formation of compounds ; it is sometimes absorbed that is, some substances have a nega- tive heat of formation ; they possess more energy than their constituents, and consequently liberate energy in decompos- 54 ing. Eeactions which liberate heat and compounds whose heat of formation is positive are said to be exothermic ; those which absorb heat and those which have a negative heat of formation are called endothermic. The measurement of the heat liberated or absorbed in chemical changes becomes therefore an important item for investigation, and many extremely interesting conclusions have been reached by the study of these phenomena. The possibility of such measurement for reactions which are,, directly and quickly realizable can be easily -understood ; and the application of a corollary of the law of persistence of energy makes it possible to determine these thermal values for reactions which can not be brought about in con- ditions suitable for heat measurement, and, in some in- stances, for reactions which can not be brought about at all. 55 This corollary as applied to chemical phenomena is known as The Law of Constant Heat Summation. It was formulated by Hess in 1840, as follows : The quantity of heat liberated or absorbed by a system of substances undergoing chemical change is dependent on the initial and final states of the sys- QUANTITATIVE LAWS OF CHEMICAL CHANGE 33 tern, and is not affected by differences in the intermediate states through which the system may pass. An illustration from the reactions which you have studied will make this clear : 16.9 grams of ammonia combine with 36.2 grams of hydrochloric acid, forming 53.1 grams of ammonium chloride, a soluble salt. Now let us suppose that we start with ammonia and hydrochloric acid as gases at the ordinary temperature, and end with the salt dissolved in water at the same temperature. The system may be passed from this initial state to this final state in two dif- ferent ways : the two gases may be directly combined, form- ing the solid salt, and then this may be dissolved in the water ; or, the gases may be separately dissolved in water and the two solutions mixed, the salt being formed in solu- tion. The law affirms that the heat disturbance, in passing from this initial to this final state, is the same, whichever method of reaching the final state is used. The law is ex- perimentally verified by^measuring the heal of solution of 16.;9*grams of ammonia gas % (,400 calories) and of 156.2 grams of hydrochloric acid gas (17.300 calories), and the heat of neutralization of the two sofufions (1#,300 calories), and summing these three quantities (=*. 38,000" calories) ; also, the heat of formation of 53.1 grams of solid ammonium chloride from the two gases (42,lJOO calories) and the heat of solution of the salt (^3,900 calories) and summing these two quantities (= 38,200 calories). The two sums are found equal within the limits of experimental error. Another method of application is seen in the determina- 55/1 tion of the heat of formation of 15.9 grams of methane from its constituents, namely, 11.9 grams of carbon and 4 grams of hydrogen; although these elements can not be made to combine directly into this compound. Methane is a combustible gas ; 15.9 grams in burning produce 43.66 grams of carbon dioxide and 35.76 grams of water ; and the heat liberated in this reaction is found to be 212,000 calories. Furthermore, 11.9 grams of carbon in burning alone liberate 34 ELEMENTARY PRINCIPLES OF CHEMISTRY 56 57 97,000 calories ; and 4 grams of hydrogen in burning alone liberate 136,800 calories. Now, starting with 11.9 grams of carbon and 4 grams of hydrogen, we may burn them sepa- rately with the requisite quantity of oxygen and convert them into 43.66 grams of carbon dioxide and 35.76 grams of water with the liberation of 97,000 -f 136,800 = 233,800 calories. Or, starting with these same quantities, we may conceive them converted into 15.9 grams of methane, the formation heat of which, a, is unknown, and then the latter substance burned with the liberation of 212,000 calories. Now, by the law, since the initial and final states are the same, x -f 212,000 = 233,800 ; hence x = 21,800 calories. One gram of hydrogen in burning to liquid water lib- erates 34,200 calories, which is more heat than is liberated by the combustion of an equal weight of any other sub- stance. The bottleful of gas which you collected in your experiments with magnesium and zinc weighed 0.2 of a gram. This, if burned, would yield enough heat to raise 2.7 such bottlefuls of water 1, or to raise 68 grams of water (nearly 2.5 ounces) from to 100. The following are the formation heats of some of the substances which have been used in your study : Water, Sulphur dioxide, Carbon dioxide, H 2 S0 2 C0 a Hydrogen sulphide, H 2 S = Mercuric iodide, Mercurous iodide, Magnesium oxide, Zinc oxide, HgI 3 Hgl MgO ZnO 68,400 calories 71,000 97,000 2,700 24,300 14,200 143,400 = 85,800 Examples of negative formation heat are : Chlorine monoxide, C1 2 Nitrogen monoxide, N 2 O Nitrogen tetroxide, N 2 () 4 Nitric oxide, NO Cyanogen, (CN) 2 Acetylene, C 2 H a 17,800 calories - 18,000 " 2,600 - 21,600 - 65,600 " - 47,600 QUANTITATIVE LAWS OF CHEMICAL CHANGE 35 Endothermic substances often show great readiness to 58 react, and are easily decomposed, sometimes explosively so e. g., chlorine monoxide and acetylene. Chemical energy and electricity. Chemical energy may 59 be transformed not only into heat, but also into electrical energy. You have observed that heat is liberated by the solution of zinc in hydrochloric acid. Now, if a sheet of zinc is partly immersed in a beaker of dilute acid, and a sheet of copper, or a carbon plate, neither being acted on by the acid, is likewise immersed in the same vessel, but separated from the zinc, and finally, if the ends FIG. 1. The galvanic cell a zinc and a copper plate (Z and C) immersed in dilute acid. not immersed be brought in contact outside the liquid, or connected by a metallic wire, the chemical energy set free by the solution of the zinc in the acid now appears, at least in part, not as heat but as the energy of the elec- tric current. This constitutes one form of the galvanic battery, in all forms of which the electrical energy is duje to the chemical changes taking place in the gener- ating cell. Indeed, it may be said that, of all the forms of energy, the one which produces chemical change is the most familiar and the most intimately associated with human activities ; at the same time it is perhaps the least understood. Not only are all those industrial applications of energy which involve the use of steam and fuel dependent on the energy of combustion, which is purely a chemical change, but, in 4 36 ELEMENTARY PRINCIPLES OF CHEMISTRY the animal organism, the heat, and the energy of muscle, nerve, and brain, are likewise derived from the chemical changes which the food and other substances undergo in the animal economy. [See Ostwald (Walker), "General Chemistry," page 209.] CLAUDE LOUIS BEKTHOLLET B. Savoy, 1748. D. Paris, 1822. (See No. 37/ 6 .) CHAPTER III COMBINING WEIGHTS-NOTATION-EQUATIONS- STOICHIOMETRY-NOMENCLATURE 1. Combining Weights THE combining weights of the elements have already 61 been defined (Nos. 43 and 44, Part I). They are multiples of the equivalent weights by a small whole number (one to five). These multiples may be regarded, for the present, as con- ventionally chosen in accordance with principles which will be explained in Chapter V. The relative active mass of an element as a constituent in any compound and as a factor in any reaction is represented as a multiple of the com- bining weight by a whole number. That this is possible is simply a consequence of the fundamental quantitative laws, studied in the last chapter. The system of com- bining weights, therefore, constitutes the basis for the quantitative expression of chemical phenomena. Later it will be seen that very important theoretical conceptions center about these values. 2. Notation The System of Chemical Notation : 62 (a) For the elements. This consists simply in represent- ing the element by a symbol usually the first letter, some- times accompanied by another, of its name ; quantitatively the symbol represents the combining mass of the element in grams. 37 38 ELEMENTARY PRINCIPLES OF CHEMISTRY (b) For compounds. This consists in representing the compound by a symbol or formula made up of the symbols of its elementary constituents, these being affected by inte- gral coefficients such that multiplication of the elementary combining weight by the coefficient shall give products which -show the relative masses of the respective elements contained in the compound. The coefficient is written below the line and follows the symbol to which it belongs. If the coefficient is one, it is omitted altogether. A coeffi- cient written on the line affects the symbols which follow it. Thus for hydrochloric acid the formula is HC1, since it contains hydrogen and chlorine as elementary constituents, and contains these in the ratio of one combining mass of hydrogen to one combining mass of chlorine that is, 1 gram to 35.2 grams. The formula of water is H 2 0, since it contains 2 grams of hydrogen to 15.9 grams of oxygen ; of ammonium chloride is NH 4 C1, since it contains nitro- gen, hydrogen, and chlorine in masses proportional respect- ively to 13.9 : 4 : 35.2. 62/1 It is clear that the formulas, H 2 C1 2 , H 4 2 , and N 2 H 8 C1 2 would represent the same relative quantities of the ele- ments as those given above, and so with any set of coeffi- cients which preserve the ratio of the first. It is custom- ary to use the simplest set of coefficients, unless there is reason for using multiples of these. The considerations which determine the choice of these possible multiples are presented in Chapter V. 62/2 The combining weight of a compound becomes therefore the sum of its elementary constituent masses, as expressed by its formula. It is sometimes called the formula weight. In the formulas of acids the H is usually placed first ; in those of bases and of salts the symbol of the metal is usually placed first; in those of oxides the is usually placed last. PROBLEMS. Calculate the relative quantities of the elementary con- stituents as expressed in the following formulas; also the combining COMBINING WEIGHTS 39 weights of the compounds : S0 2 , FeS, ZnS, H a S, FeCl 2 , FeCl s , PbCl 2 , CuS0 4 , [Al a (S0 4 ),-K,S0 4 .24H,0], C a H 4 2 , C 2 H 2 , C 6 H 6 . Reckon also the grams of the constituents contained in 100 grams of the compounds that is, the percentage composition of the latter. 3. Equations The chemical equation is an attempt to describe a 63 chemical reaction qualitatively and quantitatively in the concise form of an equation ; in it, the symbols of the fac- tors, separated by the sign +, are written on the left, and those of the products, similarly separated, on the right of the sign of equality (=), which is better interpreted by the word become or produce than by the word equal ; the sign -f- is interpreted by and or mixed with. When the symbols of the substances brought together are known, the first mem- ber is easily written ; in order to write the second, it must be learned, originally of course by observation, exactly what substances are produced when the factors are brought together ; but this depends upon conditions, as you have already had opportunity to observe, and of these the equa- tion offers no means of expression. Thus the equation Fe + S = FeS means that iron and sulphur mixed become iron sulphide, but this is not true unless the mixture be heated consider- ably above the ordinary temperature. Therefore the equa- tion, even in its qualitative aspect, is defective, and only expresses those substances which, under certain conditions, not specified, are produced from the given factors. Yet even thus limited, the equation is useful in bringing quickly to the eye the relations as to composition between the factors and the products. But the equation may be used as a means of express- 63/1 ing also quantitative relations by letting each symbol stand for a quantity of the substance named which is equal to the combining weight in grams, or in other units of 40 ELEMENTARY PRINCIPLES OF CHEMISTRY weight, and applying to it the coefficient which is neces- sary to express the relative quantities actually taking part in the reaction the active masses, as they have been called. Thus the equation Fe -f S = FeS, quantitatively interpreted, means that iron and sulphur produce iron sulphide in the proportion of 55.6 grams, or other parts by weight, of iron, to 31.8 of sulphur, to 87.4 of iron sulphide. And the equation Zn + 2HC1 = ZnCl 2 + 2H, quantitatively interpreted, means that 64.9 grams of zinc and (2 X 36.2) grams of hydrochloric acid produce 135.3 grams of zinc chloride and 2 grams of hydrogen. Clearly, then, before the equation can be used for quantitative expression, the facts as to the relative quantities must be ascertained. This can be done originally only by observa- tion, and often involves much labor and difficulty. How- ever, it may be assumed that the reaction as expressed by the equation must be in accordance with the fundamental laws of quantity just studied. Thus by Law 1 there must be represented on one side of the equation the same quan- tity of each element as upon the other side that is, the equation must " balance." Moreover, it is sometimes pos- sible, when one has become familiar with a good many reactions of the same kind, to surmise with some degree of certainty what change will take place in given condi- tions. Nevertheless, the beginning student should care- fully bear in mind that facts must be established before equations are written ; and that, because an equation is written, and written in accordance with the laws of quantity, it does not follow that it represents an actual reaction. 63/2 Again, the equation as described expresses nothing as to the energy changes which always accompany the trans- COMBINING WEIGHTS 41 formations of matter. So far as the former consist of heat changes, they are expressed by a slight extension of the ordinary chemical equations. Thus : Mg + = MgO + 143,400 calories means that 24.1 grams of magnesium combine with 15.88 grams of oxygen, producing 39.98 grams of magnesium oxide and liberating 143,400 calories of heat energy. And the equation MgO = Mg -j- 143,400 calories means that 39.98 grams of magnesium oxide, in decompos- ing into 24.1 grains of magnesium and 15.88 grams of oxy- gen, absorb the equivalent of 143,400 calories. 4. Stoichiometry It will be readily understood that when the facts as to 64 relative quantities or active masses in a given reaction are known, they may be applied, under the quantitative laws and by purely arithmetical analysis, to ascertain the actual quantities involved in specified conditions. For example : Suppose the problem is to calculate how much hydrogen would be liberated by the action of 100 grams of zinc on hydrochloric acid. The chemical fact is that it takes 64.9 grams of zinc to liberate 2 grams of hydrogen (see equa- tion in 63/j), and that the reaction always takes place in this proportion by Law 2 ; therefore by arithmetical analy- sis the proportion, 64j) : JJ : : 10.0 : # gives the desired result. Such chemical facts will be most easily recalled, probably, in equation form ; hence in solving problems of this kind 04/1 it is well first to write the equation for the reaction in- volved, then to give to this its quantitative interpretation and apply simple arithmetic. Calculation of this kind, based on the quantitative relations of reactions, is called Stoichiometry or Chemical Arithmetic. 4:2 ELEMENTARY PRINCIPLES OF CHEMISTRY 5. Nomenclature 65 Chemical nomenclature is not thoroughly systematic. Some names, so far as compounds are concerned, give, with more or less system, some indication as to the composition and the relations of the substances ; others have been arbi- trarily assigned, perhaps before the chemical nature of the substances was known. Only a few general statements are here given ; the rest may come gradually as acquaintance with substances is extended. 65/1 The elements in general combine with oxygen, and the resulting compounds are termed oxides (spelled also oxids). Examples : hydrogen oxide (commonly called water), zinc oxide, iron oxide, sulphur oxide. The ending ide (id) is generally applied to the name of one of the constituents in a compound which contains but two ; thus, carbides, nitrides, and phosphides contain respect- ively carbon, nitrogen, and phosphorus, with another ele- ment ; sulphides, chlorides, bromides, and iodides contain sul- phur, chlorine, bromine, and iodine with another element. 65/2 When there is more than one compound containing the same constituents they are distinguished sometimes by the endings ous and 'ic, as mercurous iodide (Hgl), in which the mercury shows its lower combining power, and mer- curic iodide (HgI 2 ), in which it shows its higher combining power ; and sometimes by numerals as prefixes, as in carbon monoxide (CO) and carbon dioxide (C0 2 ). 65/8 Names of acids. Acids containing no oxygen are given the prefix hydro and the ending ic. Examples : hydro- chloric, HC1, and hydrosulphuric, H 2 S. Other peculiarities are illustrated in the following series : Hypochlowus acid HC10 Chlorous acid HC10 2 Chloric acid HC10 3 Perchloric acid . . HC10 4 . COMBINING WEIGHTS 43 of salts. The salts of acids which contain no 65/4 oxygen are named by dropping the prefix hydro of the acid and changing the ending ic to ide(or id). Examples : zinc chloride, iron sulphide. The ending ic in the names of acids which contain oxygen is changed to ate in naming their salts; and the ending ous in the acid to ite in the salt. Thus the so- dium salts of the acid series above mentioned are named respectively sodium hypochlonYe, chlonYe, chlom^e, and perchlorate. Names of bases. Bases which contain hydrogen and 65/5 oxygen and another constituent (usually a metal) are now commonly named hydroxides (or ids), although they were formerly, and by some are still, called hydrates. Examples : sodium hydroxide, NaOH ; ammonium hydroxide, and zinc hydroxide, Zn (OH) 8 . STOICHIOMETRIC PROBLEMS 1. How much sulphur will exactly combine with 75 grams of 65/6 iron? 2. How much zinc with 50 grams of sulphur? 3. How much zinc is needed to make 50 grams of zinc chloride (ZnCl 2 ) ? 4. How much marble (CaCO s ) must be used with hydrochloric acid to generate enough carbon dioxide (C0 2 ) to neutralize 50 grams of sodium hydroxide (NaOH), forming sodium carbonate (Na 2 C0 3 ) f 5. How much mercury must be added to 90 grams of mercuric iodide (HgI 2 ) to convert it into mercurous iodide (Hgl)? 6. What volumes of carbon monoxide (CO) and of oxygen must be taken to produce by combination enough carbon dioxide (C0 2 ) to neu- tralize 10 grams of sodium hydroxide ? One liter of oxygen weighs 1.43 grams: one liter of carbon monoxide weighs 1.25 grams. 7. How much magnesium is needed to generate with hydrochloric acid enough hydrogen to form by burning 10 grams of water! S. Given : a sample of hydrogen containing nitrogen as impurity, and a sample of oxygen also containing nitrogen as impurity; to de- termine the percentage of impurity in each case. Measured volumes are caused to combine, forming water, liquid at ordinary temperature, ELEMENTARY PRINCIPLES OF CHEMISTRY and the following observations are made, from which calculate the impurity : RV taken = 1 vol. Hydrogen taken 2 vols.. y Hydrogen taken Oxygen taken = 10 vols. = 5 " Residual vol. after combination, RV = 3 vols. RV taken Oxygen taken = 2 vols. 1 vol. R 3 V = 3 vols. R 3 V = 2.5 vols. Hydrogen taken = 2 " R 4 V = 4.5 vols. pressure. What 9. Gas-volume = 100 c.c. at 20, and 760 mm. would this volume become at 50, and 720 mm.! 10. Calculate the percentage of elementary constituents in the com- pound whose formula is C 2 H 6 0. 11. What is the formula of the compound which contains 40.00 per cent of carbon, 6.67 per cent of hydrogen, and 53.33 per cent of oxygen f CHAPTER IV RELATION BETWEEN VOLUME, PRESSURE, AND TEMPERATURE OF GASES [THESE topics belong strictly to the subject of Physics, but inasmuch as all measurement of gas-volume involves the application of these laws, a brief study of them is here presented. If the student has already a knowledge of them from a previous course, this chapter may suitably be passed by.] 1. The Law of Boyle Relation between Volume and Pressure of Gases In a confined mass of gas, the temperature remaining con- 66 stant, the volume is inversely proportional to the pressure ; or, the product of the volume by the pressure is constant. COROLLARY. Density is directly proportional to pressure 66/2 when temperature and volume remain constant, and it is in- versely proportional to volume when temperature and pressure remain constant. The law was first discovered by Eobert Boyle about 1662. 66/3 Among Continental writers especially, it is often designated by the name of Mariotte, who, however, did not publish it until 1679. 2. The Law of Charles Relation between Volume and Temperature of Gases The volume of a mass of confined gas, the pressure remain- 67 ing constant, increases by ^ or 0.00867 of itself at for each increase of one degree in temperature ; or its pressure increases at the same rate, the volume remaining constant. 45 46 ELEMENTARY PRINCIPLES OF CHEMISTRY 67/1 COROLLARY. The laws of Boyle and Charles may he ex- pressed in one as follows : The product of volume hy pressure is proportional to the absolute temperature (that is, the observed temperature plus 273 degrees). 67/2 The law is also designated sometimes by the name of Gay-Lussac, and sometimes by that of Dalton. It was dis- covered by Charles about 1787. 67/3 NOTE I. Gaseous substances tend to depart appreciably from these laws as they approach their points of liquefaction. 67/4 NOTE II. The laws of Boyle and Charles apply to gases ' in general, independently of the chemical character of the substances. CHAPTEE V* THE RELATION BETWEEN EQUIVALENT AND COMBINING WEIGHTS AND CERTAIN SPECIFIC PROPERTIES IN studying the qualitative side of chemical phenom- 68 ena, you have been led to see somewhat of the nature of chemical change as involving the disappearance of some kinds of matter, and in their place the appearance of differ- ent kinds what has been called the change of identity in substances. The contemplation of this phenomenon gains in impressiveness with the thought that by mastering thus the production of new substances man perhaps approaches as near to independent creation as is permitted to him in any form of his material activities. Certain it is that he has made in the laboratory not only many substances which are identical with the natural products, but innu- merable others as well, which have never been found in nature. But in further studying these phenomena, particularly 69 their quantitative aspect, you learn, in addition, the per- fectly definite conditions imposed on such changes, and hence the limitation imposed on man's productiveness. Thus, no new substances can be made, save from those which contain the same elementary constituents ; that is, the con- stituent elements must be common to both factors and * If the instructor prefers to introduce at this point a part or the whole of Chapter VIII before taking up Chapters V, VI, and VII, no embarrassment will be found in so doing. The writer prefers the order herein followed. 47 48 ELEMENTARY PRINCIPLES OF CHEMISTRY products, and the change is simply a change in their dis- tribution. Nor can any kind of matter be produced with- out the disappearance of an equal mass of some other kind or kinds. 70 Summing up the laws of quantity in the definition of the equivalent and the combining weights, and accepting these for the present as determined values, we pass on to study the numerical relation between these values and others which measure very diverse properties. The first of these to be considered is the specific gravity of substances in the gaseous condition. 1. The Law of Gay-Lussac Relation between Equivalent and Combining Weights of Gases, Elementary and Compound, and their Specific Gravities 71 The specific gravity of a gas is defined as the ratio be- tween the weights of equal gas-volumes of the substance and of hydrogen, at the same temperature and pressure. This is sometimes conveniently termed the vapor-density to distinguish it from the specific gravity of the substance while in liquid or in solid condition. It is experimentally easier to weigh air than hydrogen ; therefore experimental results are often referred to air as the standard. Such values are converted to the hydrogen standard by the factor 14.40, which is the specific gravity of air referred to hydrogen as unity. In the accompanying Table II, A, No. 93, are given those of the elements which are gaseous or volatile in conditions which permit the determination of their specific gravity as gases ; with these are given also their equivalent weights, their chosen combining weights, and their vapor-densities, EQUIVALENT AND COMBINING WEIGHTS 49 as experimentally determined, hydrogen being the standard. Since the vapor-density in some cases varies greatly with the temperature of determination, and in other cases re- mains practically constant through a great range of tem- perature, some data as to temperature of determination are given ; also the boiling points. Inspection of these data shows that in the first group of 72 elements the vapor-density is numerically equal or approxi- mately equal to the equivalent weight. Now, fluorine, chlorine, bromine, and iodine, let it be recalled (see Table I, No. 49.), combine with hydrogen in the gas-volumetric ratio of 1 : 1 ; and thallium combines with chlorine in the same gas-volumetric ratio. It follows that if the equiva- lent weights in these instances be chosen as the combining weights, the formulas of these compounds for example, HC1, HBr, etc. may be interpreted as expressing the rela- tive constituent quantities, measured by gas-volume as well as by weight, since the coefficients of the elementary sym- bols give the ratio of the gas-volumes. In the second group it is seen that the vapor-density of 73 oxygen approximates very closely to equality with twice the equivalent weight. If in this case the equivalent weight 7.94 were chosen as combining weight, the formula of water would become HO, or some multiple of this, x (HO), in which the coefficients, being equal, do not give the gas- volumetric composition. If, however, the multiple of the equivalent weight by two is chosen as combining weight, the simple relation with vapor-density reappears, and the formula for water becomes H 2 0, or some multiple, x (H 2 0), in which the ratio of the coefficients, 2:1, gives the ratio of the gas-volumes. Taking thus the vapor-density as a guide in the choice 74 of the multiple of the equivalent weight which shall be used as combining weight, leads clearly in the case of tellurium to the second multiple, although the approxima- tion is not so close as in the case of oxygen. The same 50 ELEMENTARY PRINCIPLES OF CHEMISTRY rule consistently applied to selenium and to sulphur would indicate the third and the seventh multiples respectively instead of the second. These two substances, sulphur especially, show a remarkable variation in vapor-density with temperature, and it is to be noted that the lowest values approximate closely to the second multiple. How- ever, the decisive reasons for the choice of this mul- tiple must be looked for in other relations (see Ko. 87, Part I). 75 In the third group, as to nitrogen, if 4.64 be used as combining weight, the formula of ammonia must be NH, or some multiple, x (NH), which gives no indication of the gas-volumetric composition. But if the multiple by three (viz., 1&93) is chojsen, the simple relation to vapor-density NH 3 , the coefficients W3 fa^J^yjMitiojAF thej^olumes of nitrogen and hydrogen (see Table T, ivo. 49 )v But in 76 the cases of phosphorus and arsenic, the sixth multiple in- stead of the third is indicated, while for bismuth the vapor- density is nearer the second than the third, and for anti- mony it is midway between the third and fourth multiples. In all these exceptional cases decisive reasons for choice 77 must be sought in other relations. The same must be said as to the fourth group, in which zinc and cadmium and mercury show vapor-densities approximately equal to the equivalent weights, and sodium and potassium to one^ half the equivalent weights. As to the recently discovered ele- ments, helium and argon, the data are incomplete. 78 It must be borne in mind that the values for vapor- density are in most cases affected by much greater uncer- tainty than are the equivalent weights, since their deter- mination, especially at high temperatures, involves great experimental difficulty. They may, nevertheless, serve to indicate the multiple to be chosen, and their indication, in the absence of reasons for different choice, is accepted, we may say for the present, conventionally. EQUIVALENT AND COMBINING WEIGHTS 51 It is evident that to choose as combining weight for ele- 79 ments the multiple hy a whole number of the equivalent weight which approximates most closely to the vapor-density has the advantages, first, of correlating in a simple man- ner the two numerical values; and, second, in many in- stances, of bringing into one formula the expression of the proportion by gas-volume, if it is known, as well as by weight of constituents in compounds. These reasons for 80 choice of multiple receive additional weight from similar reasons which are revealed in the study of other widely different properties which will be presented in subsequent sections of this chapter. -There are also very important considerations, purely theoretic in nature, which lead to the same choice. These will be presented in connection with the atomic theory (Chapter VII). It is proper to add that the reasons of theory have had historically probably more influence, and still carry in some minds more weight, than the reasons of convenience just set forth. B In Table II, B, No. 94, are presented a few data as to the 81 vapor-density of gaseous or volatile compounds which, in- stead of being limited like the elements, are almost innu- merable. It is seen that, with the multiples which were indicated in the preceding paragraph as combining weights for the elements, the vapor-densities of the compounds ap- proximate equality uniformly with one half their combin- ing weights; whereas, with the equivalent weights thus used for the elements, the vapor-densities in the second group would approximate equality with the combining weights, and in the first group with one half the combin- ing weights, and in the third group with one and a half times the combining weights. Therefore the use of the 82 multiples chosen for the elements has still an additional advantage in bringing the vapor-density, at least of these and similar gaseous compounds, into simpler relation with 5 52 ELEMENTARY PRINCIPLES OF CHEMISTRY the combining weights of the same. And this also gives 83 basis for conventionally choosing as the combining weight of any m gaseous or volatile compound, even one whose gas-volu- metric composition is unknown, that multiple of its simplest combining weight which approximates most closely to its vapor-density multiplied by two. 84- Thus in the gas, methane, 1 gram of hydrogen is com- bined with 2.98 grams, approximately, of carbon ; but car- bon is practically non-volatile, therefore the gas-volumetric ratio is not known, and the vapor-density, at least, fur- nishes no reason to choose other value than 2.98 for the combining weight of carbon ; but the vapor-density of the compound is 8.00 ; hence we may choose the multiple by 4 that is, 16, or, more accurately, 15.9 as the combining weight of methane, and write its formula C 4 H 4 . Consider- ations of a different kind (see No. 436, Part I) have led to the choice of 11.9, the fourth multiple, as the combining weight of carbon, making the formula for methane CH 4 . 85 Again, the substance, acetylene, contains carbon and hydrogen in the ratio of 11.9 grams to 1 gram ; hence its simplest formula would be C 4 H (carbon = 2.98), or CH (carbon = 11.9), and its combining weight 12.9. But its vapor-density is 13.2 ; therefore 25.8 (12.9 X 2) is chosen as its combining weight, and C 2 H 2 is its accepted formula. 85/1 Another substance, benzene, 'likewise contains carbon and hydrogen in the ratio of 11.9 grams to 1 gram, but its vapor -density is 40. Its simplest formula would be CH, identical with the simplest one for acetylene ; but choos- ing, in accordance with this principle, the multiple by 6, gives the combining weight 77.5 and the accepted formula C 6 H 6 . 86 In the last two examples is seen still another advantage of the principle under consideration, in that it brings about a distinction in formula and combining weight between different substances of the same composition (compare No. 37/7, Part I)- (Such substances are called isomers.) EQUIVALENT AND COMBINING WEIGHTS 53 Accepting, then, this conventional rule for the choice of 87 multiple as combining weight in the case of compounds (it is more general than the corresponding rule for elements) we may reason backward to certain maximum values for the elements of exceptional vapor-density. Thus> the sub- stance hydrogen sulphide contains hydrogen and sulphur in the ratio of 1 gram of the former to 15.9 of the latter, and its vapor-density is 17.2. This leads to 33.8 grams as the combining mass, of which the constituents must be 2 grams of hydrogen and 31.8 grams of sulphur. Therefore the combining weight of sulphur can not be more than 31.8. Likewise the vapor-densities of the hy- drogen compounds of selenium, phosphorus, and arsenic indicate as maximum values for the combining weights of these elements 78.4, 30.8, and 74.44 respectively. And the vapor-densities of the chlorine compounds of antimony and bismuth indicate as maximum values for these elements 119.5 and 206.5. Effect of temperature. The simplicity of this relation 88 between vapor-density and combining weight is somewhat disturbed by the fact that, in the case of some compounds, as of some elements, the vapor-density diminishes with increasing temperature. This seems, however, to reach a limit after a certain interval of temperature. This is espe- cially notable in the cases fcf iodine, sulphur, selenium, phosphorus, and arsenic among the elements (see Table II, A), and nitrogen tetroxide and ferric chloride among the data of Table II, B. Another class of exceptional compounds contains those 89 which are also exceptions to the law of gas-volumetric pro- portions (see No. 47/ 6 , Part I). These are actually decom- posed at the temperature of observation. Upon these facts and others of the same nature is based the generalization often designated by the name of Gay- Lussac, who published it in 1808 (see also Law 5, Chap- ter II). 54 ELEMENTARY PRINCIPLES OF CHEMISTRY The law may be thus formulated : 90/1 CLAUSE I. The specific gravity of an element in gaseous condition, referred to hydrogen as unity, approximates nu- merically the equivalent weight *of the element, or a multiple of the same by a small whole number ; and it approximates equality with the combining weight. 90/2 CLAUSE II. The specific gravity of a compound in gase- ous condition (H = 1) approximates numerical equality ivith one half its combining weight. The degree of approximation as well as some of the exceptions are seen in the data of the tables. An alternative statement of the law is this : 90/3 The combining weight in grams of an element occupies, when in gaseous condition, a volume approximately equal to that of 1 gram of hydrogen at the same temperature and pressure; and the combining weight of a compound, the volume of 2 grams of hydrogen in similar conditions. 91 NOTE I. The relation between the combining weights and the specific gravities of elements in the liquid and in the solid condition is not so simple as in the gaseous condi- tion. There is, undoubtedly, a relation, but in the present state of the science it can not be satisfactorily defined. It may only be said that it is probably a periodic function of the combining weights. (The explanation of this may be deferred. See Nos. 423, 589, and 590, Part I.) 92 NOTE II. The law of vapor-densities may be deduced from the facts of gravimetric and gas-volumetric propor- tions ; for example, thus : Since 2 grams of hydrogen com- bine with 16 (using approximate values) of oxygen and form 18 of water, and since the gas-volumetric ratios are 2:1:2 respectively, it follows that 16 grams of oxygen and 9 of water must occupy a volume equal to that of one gram of hydrogen, which is the relation formulated in the law of vapor-densities. Likewise from the facts of gravimetric proportions and of vapor-densities may be deduced the law of gas-volumetric proportions. (How ?) EQUIVALENT AND COMBINING WEIGHTS 55 TABLE II, A.* The Law of Gay-Lussac 93 No. 1 8 NAME. Equiva- lent weight. Fac- tor. Combin- ing weight. Vapor- density. Tempera- ture of ob- servation. Boiling point. Hydrogen . Fluorine . . 1.0 18.91 1 1 1.0 18.91 1.0 18.23 Degrees. Degrees. Below-230 Below -95 15 Chlorine . . 35.18 1 35.18 35.83 200 -34 (i M M 23.3 1560 32 Bromine . . 79.34 1 79.34 82.77 102 59 u u " (i 52.7 1500 47 Iodine 125.89 1 125.89 127.7 253 184 " M U (i 63.7 1500 68 Thallium . . 202.6 1 202.6 206.2 1730 Red heat. 7 14 Oxygen . . . Sulphur . . . 7.94 15.915 u 2 2 u 15.88 31.83 15.90 112.0 31.8 468 1100-1719 -186 446 31 Selenium . . 39.2 2 78.4 111.0 860 665 u U K 82.2 1420 48 Tellurium . 63.25 2 126.5 130.2 1390-1439 1390 6 13 Nitrogen . . Phosphorus < 4.643 10.267 3 3 13.93 30.8 H 13.91 63.96 45.58 313 1708 -194 278 30 Arsenic. . . . 24.8 3 74.4 154.2 644 " u u " 79.5 1700 46 70 9 Antimony . Bismuth . . 39.7 68.83 3 3 119.1 206.5 141.5 146.5 1640 1640 1300 1640 Sodium . . . 22.88 1 22.88 12.7 f 742 16 Potassium . 38.82 1 38.82 18.8 f 667 27 Zinc 32.45 2 64.90 34.15 1400 930 43 Cadmium . 55.7 2 111.4 57.0 1040 770 67 73 Mercury . . Helium . . . 99.25 f 2 198.5 f 101.0 2.0 446-1730 358 Below 230 74 Argon .... f ... ! 19.8 -187 * Data mostly according to Ramsay, " Inorganic Chemistry.' 56 ELEMENTARY PRINCIPLES OF CHEMISTRY TABLE II, B.* The Law of Gay-Lussac NAME. Formula. Combining weight. Vapor-density x 2. Hydrogen fluoride HF 19 91 19 99 " chloride HC1 36.18 36.38 " bromide HBr 80 34 78 10 " iodide HI 126.89 128 Thallium chloride T1C1 237.78 237.4 Water H a O 17.88 17.98 Hydrogen sulphide . . H 2 S 33 83 34 4 " ^elenide H a Se 80 4 80 54 " telluride H 2 Te 128.5 129.6 Ammonia H 8 N 16.93 17.2 Phosphorus hydride H 8 P 33.8 33.1 Arsenic hydride H 3 As 77.4 77.8 Antimony chloride. . Bismuth " Cl 3 Sb Cl 3 Bi 224.7 312.0 224.7 327.7 Potassium iodide KI 164.7 168.9 Zinc chloride . . . ZnCl 2 135.27 132.8 Cadmium bromide Mercurous chloride CdBr, HgCl 270.2 270.1 267.0 239.6 Mercuric iodide HgI 2 450.28 468.0 Sulphur dioxide SOa 63.59 64.9 CO 27.79 27.96 " dioxide . . . C0 3 43.67 43.98 Methane CH 4 15.91 16.0 Acetylene CaHa 25.82 26.4 C 6 H 6 77.46 79.78 Nitric oxide NO 29.81 30.0 Nitrogren peroxide N0 3 45.69 45. 8 at 150 " tetroxide N 3 4 91.38 76. Oat 26 Ferric chloride . FeCl 3 161.1 155. 5 at 750-1077 u Fe a Cl 322.2 308 at 440 Data mostly according to Muir, " Principles of Chemistry." JOHN DALTON B. England, 1766. D. 1844. (See Nos. 40, 159.) EQUIVALENT AND COMBINING WEIGHTS 57 2. The Law of Dulong and Petit Relation between Equivalent and Combining Weights of Elementary Solids and their Specific Heats In the accompanying Table III, A, No. 105, are given 95 the specific heats of the solid elementary substances, together with their equivalent weights, the factors which convert the latter into combining weights, the combining weights themselves, and the products of specific heats by combining weights. Inspection shows a relation similar to that in the case of vapor-density. In some cases the equivalent weight, in other cases a simple multiple of the equivalent weight, gives numerically with specific heat a product which ap- proximates equality for all the solid elements. The prod- ucts average about 6.2 . It is seen that the approximation to constancy, save in a 96 few instances, is a remarkably close one, although the combin- ing weights range from 7 to 238, and the values for specific heat are affected with considerable uncertainty. As bearing on this, consideration must be given to the fact that specific heat varies with the allotropic condition of the substances, and with the temperature of determination. This is made evident in the data of Table III, B. Diamond, graphite, and 97 charcoal are allotropic forms of the element carbon, and the table shows how markedly the specific heats of the three forms differ, even at the same temperature of determination. It also shows that the specific heat increases with the temperature, but at a much smaller rate for high than for low temperatures. These peculiarities appear especially in those substances 98 which give exceptional products like glucinum, boron, car- bon, and silicon. As the specific heat tends to increase with the temperature, but at a rate notably diminishing as the temperature rises, it may be surmised that at a sufficient- ly high temperature the approximation to constancy would be fairly satisfactory. 58 ELEMENTARY PRINCIPLES OF CHEMISTRY 99 Eeasons similar to those suggested in considering the vapor-densities may suitably lead in the case of elementary solids to the choice for combining weight of that multiple of the equivalent weight which, when multiplied by the specific heat, gives the product approximating most closely to 6.2. This brings the combining weight into simple numerical relation with the specific heat. 100 Attention was first called to this relation by Dulong and Petit in 1819, and they urged the choice of combining weights in accordance with the above suggestion. This involved some changes in the values then in use, and their recommendation was not adopted until considerably later. The law which is generally designated by their names may be formulated thus : 101 The specific heat multiplied by the combining weight gives numerically a product which approximates a constant value, viz., 6.2, for all the solid elements. Inasmuch as this product is the quantity of heat, meas- ured, in calories (see Part II, 50/i), necessary to raise by one degree the temperature of the quantity of the substance equal to the combining weight in grams, the law may be given this Alternative Statement : 102 Quantities of the elementary solids equal to their combin- ing weights in grams have approximately equal heat capacity that is, they take the same quantity of heat, viz., 6.2 cal- ories, to effect a rise of one degree in temperature. NOTE. This relation does not hold for liquids nor for 103 COROLLARY. If the specific heat of elements in the free state remains unchanged when they have passed into combina- tion, it follows that the specific heat of a compound, multiplied by its combining weight, and divided by the number of elemen- tary combining masses ivhich it contains, should give numer- ically a value approximating 6.2 ; in other words, the heat capacity of a compound should equal the sum of the heat capacities of its constituents. EQUIVALENT AND COMBINING WEIGHTS 59 Experiment shows this relation to hold in many solid compounds which are made up of solid elements, and it is designated as the Law of Kopp and Neumann. Thus, spe- 104 cine heat multiplied by combining weight gives a product approximately 12 for the sulphides of iron and zinc, and others like them ; whereas the heat capacity of the sulphur, 5.7, plus that of the metal, 6.2 , is 11.9 . TABLE III, A* The Law of Dulong and Petit NAME. Equivalent weight. Factor. Combining weight. Specific heat. Sp. H. x comb. wt. Lithium 6.97 1 6.97 0.9408 6.6 Glucinum 4 5 2 9.0 58 5 2 Boron ... 3 62 3 10.86 0.5 (?) 5 4 Carbon (diamond). . . . Sodium 2.9775 22.88 4 1 11.91 22.88 0.459 0.2934 5.5 6.7 Magnesium 12.05 2 24.1 0.250 6.0 Aluminium 8 967 3 26 9 0.225 6.1 Silicon (crvs.) 7.05 4 28.2 0.203 5.7 Phosphorus (crys.) . . . Sulphur ; . . 10.267 15.915 3 2 30.8 31.83 0.202 0.178 6.2 5.7 Potassium 38 82 1 38.82 0.166 6.5 Calcium 19 9 2 39.7 0.170 6.8 Scandium ... 14.6 3 43.8 0.153 6.7 Chromium 17.25 3 51.74 0.100(1) 5.2 Manganese 27 285 2 54.57 0.122 6.7 Iron. 27.80 2 55.6 0.114 6.3 Nickel . 29.12 2 58.24 0.108 6.3 Cobalt 29.3 2 58.6 0.107 6.3 31.56 2 63.12 0.095 6.0 Zinc 32 45 2 64.9 0.0956 6.2 Gallium 23.13 3 69.4 0.079 (?) 5.5 Germanium 17.975 4 71.9 0.0758(?) 5.5 Arsenic (crys ) 24 8 3 74.4 0.0822 6.1 Selenium . 39 2 2 78.5 0.076 6.0 Bromine (solid) . ... 79.34 1 79.34 0.0843 6.7 Zirconium 22 425 4 89.7 0.0666 6.0 Molybdenum 47.6 2 95.3 0.0722 6.9 Ruthenium. 50.45 2 100.9 0.0611 6.2 Rhodium 51.1 2 102.2 0.058 5.9 Palladium 52.8 2 105.6 0593 6 2 Silver 107.11 1 107.11 0.057 6.1 55.7 2 111.4 0.0567 6.3 Indium . . 37.67 3 113.0 0.057 6.4 105 * Data mostly according to Muir, " Principles of Chemistry." 60 ELEMENTARY PRINCIPLES OF CHEMISTRY TABLE III, A.* The Law of Dulong and Petit (Continued) NAME. Equivalent weight. Factor. Combining weight. Specific heat. Sp. H. x comb. wt. Tin . 29 525 4 118 1 056 6.6 Antimony 39.7 3 119.1 0.0508 6.1 Iodine 125.89 1 125.89 0.0541 6.8 Tellurium 63.25 2 126 5 0474 6.0 Lanthanum 45.8 3 137.45 0449 6.1 Cerium . 34.75 4 139.0 0.0448 6.2 Tungsten 91 5 2 183 0334 6.1 Osmium 47 375 4 189.5 0.0311 5.9 Iridium 47 925 4 191.7 0.0326 6.2 Platinum 48.35 4 193.4 0.0324 6.2 Gold 195 7 1 195.7 0.0324 6.3 Mercury (solid) " (liquid) 99.25 2 198.5 0.0319 0.033 6.3 Thallium 202 6 1 202 6 0.0335 6.8 Lead 102 68 2 205.36 0.0314 6.4 Bismuth 68.83 3 206.5 0.0308 6.4 Thorium 57 725 4 230 9 0.0276 6.4 Uranium 59.45 4 237.8 0.028 6.7 106 TABLE III, B.* The Law of Dulong and Petit NAME. Temperature. Specific heat. Sp. H. xcomb. wt. Carbon, diamond 50 0.0635 0.076 < tt + 10 0,1128 1.35 tt tt 85 0.1765 2.12 it it 250 0.3026 3.63 u ft 606 0.4408 5.29 tt it 985 0.4589 5.51 Carbon graphite 50 0.1138 1.37 u u + 10 0.1604 1.93 tt U it it 61 201 0.1990 0.2966 2.39 3.56 ft tt 250 0.325 3.88 tt it 641 0.4454 5.35 tt tt Wood carbon 978 0-23 0.467 0.1653 5.60 1.95 0-99 0.1935 2.07 tt tt 0-223 0.2385 2.84 Data mostly according to Muir, " Principles of Chemistry." EQUIVALENT AND COMBININ 3. The Law of Mitscherlich Relation between Composition, and hence Combining Weight, and Specific i. e., Crystalline Form The law for present purposes may be thus stated : Substances which are similar in composition sometimes 107 show the same crystalline form. The similarity of composition here referred to is seen in the following substances : Calcium carbonate, CaC0 3 Magnesium carbonate, MgC0 3 Ferrous carbonate, FeC0 3 Zinc carbonate, ZnC0 3 They all come under the type expressed by the general formula MC0 3 , in which M represents one combining mass of the metal, and they also show the same form of crystal. The substance, sodium carbonate, Xa 2 C0 3 , on the other hand, does not have a composition similar to these, it comes under a different type ; nor does it have the same crystal- line form. The significance of this relation in its bearing on com- 108 bining weight is only of minor importance. Its application may be thus illustrated : Suppose it has been determined that the equivalent weight of zinc is 32.5, but that there is question whether to choose this or its multiple by 2 as the combining weight. With the former value, the formula for its carbonate would be Zn 2 C0 3 ; with the latter, ZnC0 3 . Suppose now that it is observed to crystallize in the same form as the carbonates of calcium, magnesium, and iron, which have the accepted symbols above given; but this crystalline form is not like that of sodium carbonate, which substance has the symbol Na 2 C0 3 . These facts would favor the choice of 65, rather than 32.5, as the combining weight for zinc. 62 ELEMENTARY PRINCIPLES OF CHEMISTRY The law was discovered by Mitscherlich and announced in 1819. 4. The Law of Raoult (I) Relation between Combining Weights of Solutes and Specific Depressions of the Freezing Points in Specified Solvent (See Part I, No. 21) 109 Preliminary statement. It has already been observed and noted (see Part I, 23/j) that the presence of substances in solution tends to lower the freezing point of the solution as compared with that of the pure solvent. 110 Law of proportionality. For the same substance in the same quantity of the same solvent, the depression of the freezing point is proportional to the quantity of the sub- stance dissolved. 110/a The specific depression of the freezing point of a sub- stance in a specified solvent is the depression produced by 1 gram of the substance dissolved in 100 grams of the solvent. 111 The law (1882). CLAUSE I. The specific depression of the freezing point multiplied by the combining weight of the solute, when the latter is compound, gives a product (D) which is approximately the same for all compound solutes in the same solvent, but differs with different solvents. 112 CLAUSE II. When the solute is elementary, its specific depression multiplied by its combining weight, or a multiple of the latter by a small whole number, gives a product (D) which is approximately the same for all elementary solutes and equal to that for compounds in the same solvent, but which differs ivith different solvents. 113 As to data. The law finds verification in many instances, but it fails in many others ; and the phenomenon is greatly complicated by the fact that the law of proportionality does not always hold good. This throws doubt on the calcula- tion of the specific depression and of the constant of de- EQUIVALENT AND COMBINING WEIGHTS 63 pression, D. Proportionality often holds while solutions are very dilute, but ceases when they become concentrated. A similar difficulty was noted in connection with vapor-density (see Nos. 88 and 89, Part I) and specific heat (see Nos. 97 and 98), the values for which vary with the temperature of determination. A very slight illustration of this is seen in the data for 114 sulphur and for iodine as solutes, which are given in Table IV, C (No. 122). The specific depression of the former varies from 0.265 to 0.248 when the concentration va- ries from 2.4 to 7.2 per cent, while that of iodine varies from 0.272 to 0.251 with a variation in per cent from 2.2 to 3.7. A great many data for compounds in various solvents 115 have been accumulated. To give some idea of the approach to constancy under the law, a few experimental values are cited in Table IV, A and B (Nos. 120 and 121). An ex- planation has been offered for such abnormally high prod- ucts as are found for the last five substances in water solu- tion (No. 121) ; it is based on the assumption that there is actual decomposition of these substances in solution (com- pare with Nos. 47/ 6 , 88, and 89). The observations for elementary substances are few, 116 since but few of the elements are soluble in ordinary solvents without chemical change. In Table IV, C (No. 122) are given the data for the constant of depression (D) in the cases of phosphorus, sulphur, bromine, and iodine. The large factors for phosphorus and sulphur should be noted in comparison with those relative to vapor-density (see Nos. 74 and 76). The constant of depression, D, for a specified solvent may 117 be determined by averaging the products for a large num- ber of substances of known combining weight, such as are exhibited in Table IV, A and B, for acetic acid and water as solvents ; but it may also be obtained by a calculation based upon the freezing point of the pure solvent and the ELEMENTARY PRINCIPLES OF CHEMISTRY latent heat of fusion, measured in calories, for 100 grams of the solvent. The calculated and the observed values agree closely in many instances. A few examples are given in Table IV, D (No. 123). 118 An alternative statement of Eaoult's law may be made as follows : Masses of solutes which are proportional to their combining weights, or in the case of elements to simple mul- tiples thereof, when dissolved in 100 grams of the same sol- vent, produce approximately equal effects in depressing the freezing point of the solutions as compared with that of the solvent. 119 Method of applying. Suppose it is questioned what multiple of 16.9 shall be used as the combining weight of hydrosulphuric acid. Experiment gives 1.05 as the specific depression for this substance in acetic acid. Hence, by Eaoult's law, OQ Comb. wt. = -^ = 37; J..UO and the multiple of 16.9 by 2 is chosen, since it approximates most closely to 37 . 120 TABLE IV, K*Raoults Law (I) Solvent, acetic acid. Freezing point = 16.75. Constant, D = 39. SOLUTE. Formula. Product, D. Alcohol C 2 H 6 36.4 Ether C 4 H 10 39.4 Chloroform .... CC1 3 H 38.6 Glycerine C 8 H e 3 36.2 Naphthalene 39.2 Camphor Ci H 16 36.4 Water H 2 33.0 Carbon disulphide CS 2 38.4 Hydrosulphuric acid .... H 3 S 35.6 Sulphur dioxide S0 2 38.5 Sulphuric acid H 2 S0 4 18.6 Hydrochloric acid . HC1 17.2 Mg(C 2 H 8 9 )3 18.2 JEREMIAS BENJAMIN RICHTEE B. Germany, 1762. D. 1807. (See No. 41/ 9 .) EQUIVALENT AND COMBINING WEIGHTS 65 TABLE IV, B.* Raoulfs Law (I) Solvent, water. Freezing point, 0. D = 19. 121 SOLUTE. Formula. Product, D. Alcohol C 2 H 6 17.3 Glycerine . . ... C 3 H 8 3 17.1 Cane sugar 18.5 C 2 H 4 O 2 19.0 Magnesium sulphate MgS0 4 19 2 Ferrous " ... FeS0 4 18 4 Zinc " ZnS0 4 18.2 Copper CuS0 4 18.0 Hydrochloric acid. HC1 39 1 Ammonium chloride NH 4 C1 34.8 Potassium " KC1 33.6 " tartrate K 2 C 4 H 4 6 - 36.3 Pb(N0 3 ) 2 37.4 TABLE IV, C.Raoulfs Law (I) 122 Concen- SOLUTE. Solvent. tration. Grams in 100 of Spec. D. Multiple of comb. wt. D (exp.). D (cal.). solvent. Phosphorus f . Sulphur f Benzene .... Naphthalene. 0.72 2.42 0.392 0.265 30.8x4 31.8x8 48.3 67.4 51.0 69.4 " 7.20 0.248 Bromine \ , . . . Acetic acid. . 1.71 0.251 79.3x2 39.8 38.8 Iodine \ Naphthalene. 2 19 0.272 125.9x2 68.5 69 4 <* 3.72 0.251 TABLE IV, D.* Raoulfs Law (I) 123 SOLVENT. D observed average. D calculated. Water 18.5 18.9 Acetic acid . ... 39.0 38.8 Benzene 49.0 51.0 Naphthalene . . 71.0 69.4 * According to Ostwald (Muir), " Solutions." f According to Hertz, Zeitschrift fur physikalische Chemie, VI, 358 (1890). According to Paterno and Nasini, Berichte der deutschen che- mischen Gesellschaft, xxi, 2153 (1888). * According to Nernst (Palmer), " Theoretic Chemistry." 66 ELEMENTARY PRINCIPLES OF CHEMISTRY 5. The Law of Baoult (II) Relation between Combining Weights of Solutes and Spe- cific Elevations of the Boiling Temperature in Specified Solvent 124 Preliminary statement. The facts presented in the pre- ceding section relative to the freezing point of solutions are closely paralleled by the facts relative to the boiling temperature of the same. Attention has already been given to the fact that the presence in solution of iion-vola- tile substances, or of substances practically non-volatile in the actual conditions tends to raise the temperature of the boiling solution as compared with that of the pure solvent (compare No. 24/ 4 ). This temperature is not to be con- fused with that of the boiling point, which is strictly the temperature of the vapor, not of the liquid. The term "boiling temperature" is herein used to designate the first. 125 Boiling takes place when the pressure of the vapor which is formed in the body of the liquid just exceeds the pressure of the atmosphere plus the resistance which the liquid offers to the formation and movement of bubbles. Now, the presence of non-volatile substances in the liquid reduces the vapor-pressure for a given temperature, and consequently increases the temperature which is necessary to produce a vapor-pressure equal to the opposing pressure of the atmosphere ; in other words, it raises the boiling temperature of the liquid, although the temperature of the vapor, after it has escaped from the liquid i. e., the boiling point may be unaffected. 126 The law of proportionality. For the same substance in the same quantity of the same solvent the elevation of the boiling temperature is proportional to the quantity of the substance dissolved. The specific elevation of the boiling temperature of a EQUIVALENT AND COMBINING WEIGHTS 67 substance in a specified solvent is the elevation produced by 1 gram of the substance dissolved in 100 grams of the solvent. The law (1886). CLAUSE I. The specific elevation of 127 the boiling temperature multiplied by the combining weight of the solute, when the latter is compound, gives a product (E) which is approximately the same for all compound solutes in the same solvent, but differs with different sol- vents. CLAUSE II. When the solute is elementary, its specific 128 elevation multiplied by its combining weight, or a multiple of the latter by a small whoU number, gives a product (E) which is approximately the same for all elementary solutes and equal to that for compounds in the same solvent, but which differs with different solvents. As to data. The statement made concerning exceptions 129 both to the law of proportionality and to the law of Kaoult relative to freezing point is equally applicable to the phe- nomenon of boiling temperature. There are many verifica- tions, but also many exceptions. Variation from propor- tionality is seen in the data for sulphur, Table V, F. For a concentration of 1.5 per cent the multiple of the combining weight by 8 gives a good approximation to the expected product ; whereas at a concentration of 10 per cent the factor 9 gives the closest approximation. A few of the very many experimental determinations for compounds are given in Table V, E, and those of the available elements in Table V, F. The constant of elevation, E, for a specified solvent may 130 also be calculated from the experimentally determined boiling temperature of the solvent and the latent heat of vaporization, measured in calories, for 100 grams of the same. An alternative statement of this law may be made in 131 terms exactly similar to those employed in the law relative to the freezing point of solutions. 68 ELEMENTARY PRINCIPLES OF CHEMISTRY TABLE V, E.* Raoulfs Law (II) 132 Solvent, alcohol. Boiling point =78. 3. Constant, E= 11. 5 (calculated). SOLUTE. Formula. Comb. wt. Products. Concen- tration. Spec. E. Naphthalene Ci H 8 CN 2 H 4 C 7 H 6 2 HgCl 2 LiCl KC 2 H 3 O a ing point, 10( 128 60 122 271 42.5 98 ). Cons 10.0 11.1 12.2 11.7 13.2 12.0 bant, E = 5.2 (cal culated). Urea Benzoic acid Mercuric chloride. . . Lithium " Potassium acetate. . . 133 Solvent, water. Boil Cane sugar Ci 2 H 22 Oii C 6 H 14 6 CN 2 H 4 H 3 B0 3 HgCl a CdI 2 NaC 2 H 3 0, 342 182 60 62 271 366 82 4.9 5.0 4.3 4.8 5.0 5.2 9.6 Mannite . . . . Urea Boric acid Mercuric chloride. . . Cadmium iodide. . . . Sodium acetate. . 134 Solvent, carb. disulphide. Boiling point, 46.2. Constant, E = 23.7 (calc.). Anthracene C 14 Hio 178 24.4 0.59 0.137 23.7 4.02 0.132 Naphthalene C 10 H 8 128 23.2 2.99 0.181 21.9 7.54 0.171 Camphor C, H, 8 152 21.4 2.02 0.141 u 21,1 8.09 0.139 TABLE V, F.f Raoulfs Law (II) 135 Solvent, carbon disulphide, Boiling point, 46.2. Constant, E = 23.7. SOLUTE. Concentration. Per cent. Spec. E. Multipleof comb. wt. Product, E. Phosphorus u Sulphur 1.58 10.80 1.52 10.00 0.185 0.159 0.094 0.083 30.8x4 a 31.8x8 c - 22.8 19.6 24.0 21.0 Iodine 1.27 0.095 125.9x2 23.9 9.00 0.087 it 21.9 * According to Beckmann, Zeitschrift fur physikalische Chemie, vi, 437 (1890). t Ibid., v, 76 (1890). EQUIVALENT AND COMBINING WEIGHTS 69 SUMMAET In summing up the facts presented in this chapter rela- 136 tive to the elements, it is seen that masses of these sub- stances, which are proportional to their equivalent weights or to products thereof by a small whole number, produce equal, or approximately equal, effects in occupying volume when in the gaseous condition ; in absorbing heat with equal changes of temperature when in the solid condition ; and in depressing the freezing point and raising the boiling temperature when in dilute solution by equal quantities of the same solvent. Moreover, it is learned that simplicity of relation is 187 brought about by choosing for combining weight the exact multiple of the equivalent weight which in the gaseous state occupies a volume most closely approximating that occupied by 1 gram of hydrogen in similar conditions of temperature and pressure, and which in the solid state requires a quantity of heat most closely approximating 6.2 calories to effect a rise of one degree in tempera- ture; and that the choice is conventionally made in ac- cordance with these principles, unless reason is found for choosing otherwise. Relative to compounds, it is learned that masses of these 138 equal to their simplest combining weights in grams do in many instances occupy in gaseous condition approxi- mately equal volumes, and that this volume is equal to that occupied by 2 grams of hydrogen at the same tempera- ture and pressure. And again, for simplicity of relation, that exact multiple 'of the simplest possible combining weight is conventionally chosen for the actual combining weight which in the form of gas occupies a volume most closely approximating that of 2 grams of hydrogen. In addition, it is learned that masses of compounds which are proportional to their combining weights likewise in many instances produce equal effects in depressing the freezing TO ELEMENTARY PRINCIPLES OF CHEMISTRY point and in elevating the boiling temperature when in solution under similar conditions. 139 And, finally, in comparing elements with compounds it is to be particularly noted that the combining masses of the elements must generally be multiplied by a small factor in order to produce equal eifects with those of the combin- ing masses of the compounds, in respect to the specific properties herein studied. In the matter of vapor-density (see Table II, Nos. 93 and 94) the factor is generally 2. Exceptions may be noted as follows : for iodine the factor is 2 at low temperature and 1 at high, but it is 2 in relation to freezing point and boiling temperature ; for sulphur the factor is 7 at low temperature, 2 at high, and 8 or more in relation to freezing point and boiling temperature ; for phosphorus it is 4 at low temperature and the same for the phenomena of the freezing point and boiling temperature ; for arsenic, relative to vapor-density, it is 4 at low and 2 at high temperature, while it is 1 for sodium, potassium, zinc, cadmium, and mercury. 140 Relative to specific heat, on the other hand, the factor varies and is equal to the number of elementary combining masses which the compound contains. 141 Great significance is attached to the facts of this chap- ter, and important theoretic conceptions are based upon them. The correlation of these widely diverse phenomena, and there are likewise others which have not been here presented, with those of chemical change can not but be impressive to the thoughtful student of nature. It forces upon him the conviction that there must be some cause underlying the common relation ; and men have not been slow in their attempts to discover and explain this cause. The outcome of their efforts will be set forth in connection with the atomic theory in Chapter VII ; but, preceding this, it is well to give brief consideration to the method of deter- mining experimentally these fundamental values, the equiva- lent and combining weights. CHAPTER VI METHOD OF DETERMINING EQUIVALENT AND COMBINING WEIGHTS OF ELEMENTS AND FORMULAS OF COMPOUNDS 1. Determination of the Equivalent Weight of an Element THE equivalent weight of an element has already been 142 defined (see Nos. 42 and 43). The most direct method of determining its value is to ascertain by experiment what mass of the substance combines with 1 gram of hydro- gen. If it combines in more than one ratio with hydrogen, under the law of multiples, more than one value will be found. For convenience the smallest might be designated, although this is not important. If it forms no compound with hydrogen, it is necessary to determine its ratio of com- bination with some element, the equivalent weight of which has been determined. Many of the elements do not com- bine with hydrogen, or, if they do so combine, they form compounds not so well suited for analysis as those with some other element. On the other hand, most of the elements do combine with oxygen and furnish compounds 143 available for analysis. With chlorine, also, many advan- tageous compounds are formed. And so, historically, the oxides and chlorides have been more frequently used than the hydrides, and in some instances other and more com- plicated compounds than these have served. To determine the equivalent weight, therefore, it is ne- 144 cessary to ascertain by experiment ivhat mass of the element combines ivith 1 gram of hydrogen, or, if this is impracti- cable, with 7.94 grams of oxygen, or 85.189 grams of chlorine, or with the equivalent mass of some other substance. 71 72 ELEMENTARY PRINCIPLES OF CHEMISTRY 145 There is no need to consider here the details of experi- mental procedure. The most refined analytical methods known and available to chemists have been brought to bear upon the determination of these fundamental values. Not by any means the same degree of accuracy has been at- tained in all instances. The values given for the combining weights in Table XI, No. 644, have in most instances an uncertainty of one or two units in the last figure. The greatest accuracy is claimed for that of silver, with an un- certainty of less than four units in more than one hundred thousand, an accuracy " which has scarcely been obtained elsewhere in the exact sciences, much less surpassed." * The name of Stas (1860-1865), a Belgian chemist, is mem- orably associated with this great work. 146 Several other values show an uncertainty of only a few units in ten thousand or more, while in others it is as much as one per cent. The combining ratios of oxygen and chlorine with hydrogen are especially important, being the values upon which many of the others depend, and vast labor and great skill have been brought to bear upon their determination. In spite of this, the equivalent of oxygen, until a comparatively recent date, has been affected by an uncertainty as great as one part in two hundred. For this reason it has been proposed to fix this value as eight that is. to make 8 grams of oxygen the base of the sys- tem instead of 1 gram of hydrogen, letting the uncertainty lie upon the value for hydrogen, the equivalent weight of which would become one and a small fraction. The com- bining weights reckoned on this basis (0 = 16) are also given in Table XI, No. 644. 147 By experiments involving the greatest skill and most elaborate apparatus, Morley, of Cleveland (1895), and others, have determined the mass of oxygen which combines with 2 grams of hydrogen as 15.879, with a probable error of * Ostwald (Walker). " Outlines of General Chemistry." DETERMINING EQUIVALENT WEIGHTS 73 only a few units in about sixteen thousand. A few of Morley's results are here quoted as of interest : 15.877 15.877 15.877 15.878 15.879 15.881 15.881 15.882 15.8792 = the mean of these and others. As of special interest, also, in connection with your ex- 148 periment under Law 4, Chapter II (No. 41/ 2 ), are quoted some recent results in determining with the utmost exact- ness the mass of zinc which combines with 16 grams of oxygen, by converting a weighed quantity into nitrate, igniting, and weighing the oxide obtained : 65.459 65.445 65.459 65.440 65.489 65.475 65.437 65.447 65.457 * In the case of some of the metals, the experimental de- 149 termination of the ratio of hydrogen displacement, as illus- trated in your experiments (Nos. 41/ 3 and 41/ 4 ), is useful, but rather as a secondary method of control than as a direct determination of the equivalent weight. * Morse and Arbuckle, American Chemical Journal, xx, 195 (1898). 74 ELEMENTARY PRINCIPLES OF CHEMISTRY 2. Determination of the Combining Weight of an Element 150 The choice of that multiple of the equivalent weight which shall be designated as the combining weight is con- trolled by such one or several of the relations given in Chapter V as may be applicable to the element in question. In the event of conflicting indications, the law of Gay-Lus- sac probably carries the greatest weight, and next to this the law of Dulong and Petit. The laws of Eaoult have come into use in only comparatively recent years, and have given valuable indications in many cases in which the data for the application of the other laws were lacking. 151 Besides the reasons for the choice of multiples, which are found in the correlation of the specific properties studied in Chapter V, a very important aid is obtained in considering the relation of the element to other elements through its general properties. This relation is embodied in the law of periodicity (see Chapter VIII, Nos. 436 and 589, Part I), the presentation of which is best deferred until after a somewhat detailed descriptive study of the elements themselves. Suffice it for the present to state, by way of illustration, that the relation of carbon to the other elements, now expressed by the law of periodicity as de- pendent on the combining weight 12, would be quite inex- pressible as dependent on any other multiple of its equiva- lent weight 3, if the other elements retained their present values. 152 It may be well to emphasize by repetition here what is implied in all the foregoing study, that these relations with specific properties are used only to determine the choice of that exact multiple by a whole number of the equivalent weight which shall be used as combining weight, and not to determine the fundamental ratio itself, which is expressed in the equivalent weight. This value is determined by chemical analysis with a precision which, save in a few in- JOSEPH LOUIS GAY-LUSSAC B. France, 1778. D. Paris, 1850. (See Nos. 47/ T , 89, 220.) DETERMINING EQUIVALENT WEIGHTS f5 stances, far exceeds that reached in the measurement of the specific properties. The approximations in the latter are, however, usually sufficient to distinguish clearly between different multiples. 3. Determination of the Formula of a Compound The first step in the solution of this problem must be to 153 determine qualitatively what elementary constituents the given compound contains. It suffices, and may be more convenient, to learn the proximate constituents, if the com- position of the latter is known. Thus carbon dioxide may be recognized as a constituent of sodium carbonate (com- pare Exps. Nos. 35 and 81/i), and by this it is known that carbon and oxygen are among its elementary constituents. It is often more practicable to convert an elementary con- stituent into some recognizable compound, than to separate and identify it in its elementary condition. Thus, if a compound is combustible, and water is identified as a prod- uct of its combustion, the presence of hydrogen as a con- stituent is proved (compare Exps. 34/ 2 and 34/ 4 ). The second step must be to ascertain, as accurately as 154 possible, the quantities of the elements present in a meas- ured quantity of the compound. And here likewise, it may be necessary to convert the elements quantitatively into forms of known composition, suitable for measurement. The results of such quantitative determination are usually expressed as parts found in one hundred parts of the com- pound taken. To deduce a formula when the percentage composition 155 is known, divide the percentage of each constituent by the combining weight of the same. A series of quotients is thus obtained which should stand to each other in a ratio closely approximating that of whole numbers ; if the whole num- bers desired do not appear on inspection, divide the quo- tients by the smallest of the series. Write the symbols of 76 ELEMENTARY PRINCIPLES OF CHEMISTRY the elementary constituents together to form the symbol of the compound, and attach to them, as coefficients (written below the line), the simplest set of whole numbers which stand in the ratio of the quotients. This formula shows the percentage composition of the substance, and the sum of the constituent weights becomes its combining weight, unless there be found some reason for choosing a multiple of this formula and this combining weight, as indicated by the law of Gay-Lussac, or those of Eaoult. Of these the former is reckoned as giving the surer indication. There are many compounds, however, to which neither can be applied. 155/1 The following example of actual experimental results is taken from a recent journal : * Experiment I. Experiment II. Carbon 12.00 per cent 11.69 per cent Hydrogen 3.98 " 4.02 Nitrogen 6.77 Bromine. . . 77.56 " 100.31 12.00 -5- 11.9 = 1.008 3.98 -f- 1.0 = 3.98 6.77 -f- 13.9 = 0.487 77.56 -s- 79.3 = 0.978 1.008 -^ 0.487 = 2.07 3.98 -=-0.487 = 8.17 0.487 -=- 0.487 = 1.0 1.978 -r- 0.487 = 2.01 Hence is deduced the formula C 2 H 8 NYBr2, which expresses the follow- ing percentages to which the experimental values approximate : Carbon =11.65 Hydrogen = 3.92 Nitrogen = 6.80 Bromine = 77.63 100.00 156 EXAMPLES. (1) Deduce the formula of the compound from the following data : 92.30 per cent of carbon, 7.70 per cent of hydrogen ; specific gravity as a gas (H = 1) is approximately 38. * American Chemical Journal, vol. xx, p. 56 (January, 1898). DETERMINING EQUIVALENT WEIGHTS Y? (2) Also for 40.00 per cent of carbon, 6.67 per cent of hydrogen, and 53.33 per cent of oxygen ; specific gravity, as a gas, is approxi- mately 29. (3) Also for 78.86 per cent carbon, 10.60 per cent hydrogen, 10.53 per cent oxygen. (4) Assume that 1 gram of a substance (e. g., alcohol) containing only carbon, hydrogen, and oxygen yields by burning 1.913 grams of carbon dioxide and 1.173 grams of water (the quantity of oxygen is assumed to be the difference between the total and the sum of the car- bon and the hydrogen) ; assume that the specific gravity of the gas is approximately 23 ; deduce the percentage of carbon, hydrogen, and oxygen, and the formula of the substance. (5) Experiment gives for ferric chloride the vapor-density 10.7 (air = 1) and the percentage of chlorine 65.76 ; also the specific heat of iron as 0.114; assume for the combining weight of chlorine 35.2 and deduce the combining weight of iron and the formula of iron chloride (Meyer and Griinwald). (6) What percentage of chlorine is contained in ammonium chloride? (7) What is the percentage of iodine in mercurous iodide? In mer- curic iodide t (8) Assume that 10 grams of pure iron displace 0.36 of a gram of hydrogen from hydrochloric acid ; calculate the equivalent weight of iron. (9) What weight of oxygen should be obtained by heating 10 grams of pure potassium chlorate ? (10) What is the percentage of carbon dioxide contained in calcium carbonate ? * CHAPTER VII* . THE ATOMIC THEORY 157 IT is the purpose of this chapter to present under this general title the prevailing theories which have been de- vised to explain the facts already set forth, as well as many other facts impracticable of treatment in an elemen- tary study. 158 Some writers on the history of chemistry endeavor to trace the origin of the atomic theory back to Greek and Latin speculative writers, who did indeed discuss the nature of matter, its infinite divisibility, and kindred subjects; but this was rather metaphysical speculation, and may hardly be regarded as related to physical science in the strict meaning of the term. 159 To Dalton, a chemist and physicist of Manchester, Eng- land, is credited the invention of the atomic theory, since he was the first to give it quantitative form, and to make it a truly scientific hypothesis, to be tested by experiment and observation. This he did between 1803 and 1806, and he was led to the conception largely by his discovery of the fact of multiple proportions. Since then the theory has been greatly modified by change and extension, but still preserves his fundamental idea. 160 In order to present the theory as a whole, and logically * If the instructor desires to defer the presentation of the Atomic Theory until later in the course, there is nothing in the arrangement of the subject-matter to hinder so doing. 78 THE ATOMIC THEORY 79 rather than chronologically, it is necessary to give atten- tion, first, to a purely physical hypothesis as to the consti- tution of matter. The assumption which most satisfac- torily explains many facts in the domain of physics is that gross matter that is, matter as it appears to the senses is an aggregation of very small material particles, separated by intervening spaces, and that these particles are the units (or individuals, so to speak) upon which act different kinds of energy, such as heat, light, and electricity. The supposition is that substances (for example, glass), although they appear homogeneous, are in reality grained, and would so appear to a sufficiently magnified sense as a pile of shot, -at a distance, would seem homogeneous, but, seen closely, would show its grained structure. These particles are called molecules, and the theory is 101 named the molecular theory. Upon it depends largely the explanation of the important phenomena of light and of heat, such as expansion and contraction with change of temperature, the solid, liquid, and gaseous states, and, in respect to the latter, the law of Boyle and that of Charles. These molecules are reckoned as real magnitudes, fur- 162 nishing as definite a basis for mathematical reasoning as if they could be individually weighed on the balance. They are much smaller than anything revealed by the most pow- erful microscope, yet they are capable of approximate meas- urement by indirect calculation. Lord Kelvin, the eminent English physicist, has estimated that if a drop of water were magnified to the size of the earth, and its molecules in the same proportion, the mass would appear more coarsely grained than a heap of small shot, but less coarsely grained than a heap of baseballs. Or again, a cube one four-thou- sandth of a millimeter or one one-hundred-thousandth of an inch on the edge is about the smallest mass to be seen by a good microscope, and this contains from sixty millions to one hundred millions of molecules. Such magnitudes are 80 ELEMENTARY PRINCIPLES OF CHEMISTRY so far beyond the range of ordinary experience that they seem as inconceivably small as the celestial dimensions seem inconceivably large. The atomic theory (or the atomic-molecular theory, as it might be called) in its present form involves the follow- ing fundamental assumptions : 163 1. The molecular constitution of matter. It is assumed that gross matter is an aggregation of very minute mate- rial particles, separated by intervening spaces. These par- ticles, called molecules, act as units to all forces other than chemical. 164 2. The kinetic theory of gases (due largely to Clausius, 1857) assumes that in the gaseous condition the separating spaces are considerable as compared with the size of the molecules ; that the molecules are in rapid motion in all directions, colliding constantly with each other and with the walls of the containing vessel, and, being perfectly elastic, rebounding after every collision ; and that gas pressure is due to this impact on the walls. From these assumptions, and by the application of the laws of me- chanics, may be deduced the laws of Boyle and Charles, which also have a purely experimental basis, as has been already described. 165 3. Avogadro's hypothesis. It is assumed that equal vol- umes of all gases, independently of their chemical char- acter, at the same temperature and pressure, contain the same number of molecules. This was put forth as a hy- pothesis by Avogadro in 1811, but it may also be deduced from the kinetic theory of gases which was developed con- siderably later. 166 4. The chemical definition of a molecule defines it as the smallest mass of a substance in which the properties of the substance inhere; that is, the identity of a substance is conceived as resident in its molecule. It is assumed that all molecules of the same substance are alike and have the same mass. THE ATOMIC THEORY 81 5. As to molecular weights. It is assumed that the rela- 167 tive mass of molecules of different substances, referred to the molecule of hydrogen as unity, is equal to the specific gravity of the substance in gaseous condition, referred to hydrogen as standard. It follows that the values which have been called the combining weights of compounds are by this theory called their molecular weights (in this sense not expressible in grams), and that the molecular weight of hydrogen, and of some other elementary gases, is twice the combining, weight. This assumption is a direct deduction from the laws of Gay-Lussac (see Part I, Nos. 47, 90/ l5 and 90/ 2 ) and the hypothesis of Avogadro. 6. The divisibility of the molecule of a compound into 168 smaller parts which are unlike each other, and unlike the original molecule, is assumed because a measured mass of the compound can be separated into smaller masses of its elementary constituents. Thus, since 18 grams of water can be separated into 2 grams of hydrogen and 16 grams of oxygen, it is assumed that the molecule of water, weigh- ing nine times as much as the molecule, of hydrogen, can be, and in the decomposition of water is, separated into smaller parts of two different kinds, called atoms, one in the aggregate showing the properties of hydrogen, and the other those of oxygen. Compare No. 172. 7. The divisibility of the molecule of an element, at least 169 of some elements, is assumed, and the assumption is based on reasoning, of which the following is an example : It is a fact that one gas-volume of hydrogen and one gas-volume of chlorine combine and form two gas-volumes of hydro- chloric acid (see Part I, No. 49/j). It is assumed that two gas-volumes of the compound contain twice as many mole- cules as one gas-volume of hydrogen and twice as many as one gas-volume of chlorine. It is also assumed that every molecule of hydrochloric acid contains its smaller particle of hydrogen and of chlorine. Therefore it is assumed that every molecule of hydrogen, likewise of chlorine, is divisible 82 ELEMENTARY PRINCIPLES OF CHEMISTRY into at least two equal and like parts, since the hydrogen molecules are distributed among twice their number of molecules of hydrochloric acid. Exactly similar reasoning leads to the assumption that the molecules of oxygen in forming water, and those of nitrogen in forming ammonia, are divided into at least two equal and like parts. And the same is true of most of the elementary gases. These smaller particles are named atoms. 170 This assumption in distinguishing between the mole- cule and the atom of elements interprets the facts concern- ing vapor-density (see Nos. 72, etc., 93, and 94) as indicating that the molecule of hydrogen and that of most of the elementary gases contain two atoms ; that the molecule of iodine contains two atoms at low temperature, but only one at high in other words, the atom and molecule become the same ; likewise the molecule of sulphur contains seven or eight atoms at low and two at high temperature ; that of phosphorus and that of arsenic contain four atoms ; while the molecules of sodium, potassium, zinc, cadmium, and mercury contain Jmt one atom. See Nos. 139 and 141. 171 8. The atom is defined as the smallest mass of each ele- mentary substance that is found in any molecule; it is the unit upon which chemical force acts, remaining undivided through all changes. It is assumed that all atoms of the same element are alike, but unlike the atoms of every other element. 172 It is customary to speak of the atoms of hydrogen and of the atoms of oxygen, yet we may not be justified in assuming that an aggregation of such atoms, uncombined in molecules, would show the properties of hydrogen and of oxygen. 178 9. As to atomic weights. It is assumed that all chem- ical changes are due to the interaction of atoms; and that the relative mass of atoms of different elements, referred to the atom of hydrogen as unity, is constant and numerically equal to the combining weight of the element. Therefore THE ATOMIC THEORY 83 the values which have been called combining weights of the elements are by this theory called the atomic weights (in this sense not expressible in grams). It follows that the formula of a compound shows the kind and number of atoms in its molecule. This assumption constituted the atomic theory as first 174 announced by Dalton in 1804. It is simply a theoretic interpretation, or explanation, of the facts embodied in the laws of fixed, multiple, and equivalent proportions. 10. As to heat capacity. It is assumed that all atoms 175 have the same heat capacity (see Nos. 101 and 102). This is but the theoretic interpretation of the law of Dulong and Petit. 11. As to Raoult's laws. The theoretic statement is that 176 the effect of a solute in lowering the freezing point and in raising the boiling temperature of a solvent is dependent on the number and not on the kind of molecules of the solute present in a specified mass of the solvent. The observations concerning sulphur and phosphorus and iodine in solution are interpreted as indicating eight, four, and two atoms respectively in the molecule (compare Nos. 116, 122, 135, 139, and 141). The exceptional values for com- pounds, such as those seen in the last three items of the table, No. 120, are thought by some to indicate, for the substances when dissolved in the specified solvent, a molec- ular weight which is double the formula weight given in the table. On the other hand, values such as those seen in the last five items of No. 121 and in the last item of No. 133 are interpreted as indicating, not that the molecular weights of the substances in question are one half those assigned in the table, but that the substances are actually decomposed in the conditions of observation into elementary or into proximate constituents, so that there are twice as many molecules present as there would be without decomposi- tion. This assumption as to the peculiar condition of some substances when dissolved in some solvents is the basis of 7 84 ELEMENTARY PRINCIPLES OF CHEMISTRY what is known as the theory of electrolytic dissociation or of ionization. It was advanced by Arrhenius in 1887, and it includes the theoretic interpretation of many of the phe- nomena of solution besides those pertaining to the freezing point and to the boiling temperature. 177 12. As to the structure of molecules. It is assumed that the properties of substances are affected not only by the kind and number of atoms in the molecule, but also by their arrangement, grouping, or linkage. 178 The discovery in 1828 by Wohler that two different substances may have the same percentage composition found no explanation in the theory of that time, which assumed that the properties of the compound were de- pendent only on the kind and number of its constituent atoms. Since then very many instances have been revealed of substances which are identical in percentage composi- tion, but still very different in properties (compare ^os. 37/7, 85, 85/j, and 86). This has led to the assumption of molecular structure to explain the existence of such substances, and many of the most conspicuous achieve- ments of modern chemistry may be properly regarded as the outcome of this conception, or of experiments guided by it. 179 Substances having the same percentage composition, but not identical, are called isomers. Of these there are two varieties the polymers and the metamers. The poly- mers are substances which have the same percentage com- position, but differ in molecular weight. For example : the two substances acetylene, C 2 H 2 , and benzene, C 6 H 6 , are polymers. In terms of the theory, the molecule of the first contains two atoms of carbon and two of hydrogen, while that of the second contains six atoms of carbon and six of hydrogen. 180 The metamers are substances which have the same percentage composition and the same molecular weight. Their difference of properties is theoretically explained as THE ATOMIC THEORY 85 due to difference in the grouping of the constituent atoms of the molecule. This is expressed in their formulas by a difference in the grouping of the symbols. For example : the first instance of isomerism, discovered by Wohler, was in the two substances ammonium cyanate, a salt, NH 4 CNO, and urea (NH 2 ) 2 CO, a very different substance having not even the general characteristics of the salts. The differ- ence in structure is shown in the different arrangement or grouping of the elementary symbols. Evidence as to structure. The subject of structure finds 181 its greatest development in the study of the carbon com- pounds, usually called organic chemistry, and no detailed consideration of it is judged suitable for a course having the scope of this one. It is desired, however, to give a suggestion concerning the kind of evidence upon which assumptions as to structure are based. For illustration, take the substance known as acetic acid. It contains the elements carbon, hydrogen, and oxygen. The percentage of these gives the formula CH 2 0, the combining weight of which would be 30. But the specific gravity of the sub- stance in gaseous condition indicates the combining weight 60; therefore the molecular formula of the substance is C 2 H 4 2 . This is polymeric with another very different substance which has the formula CH 2 0. Now, it is ob- served that the substance C 2 H 4 2 acts as an acid that is, it contains hydrogen which can be replaced by the action of metals, forming a series of well defined salts. But ex- periment shows also that only one fourth of the hydrogen contained can be thus replaced. The theoretic interpreta- tion of this is, that one of the four hydrogen atoms is held or linked in the molecule in a manner somehow differing from that of the other three, and this, the first step in differentiating the atoms in the structure of the molecule, is expressed by writing the formula H(C 2 H 3 2 ). Again, experiment shows that by acting upon acetic 181/1 acid with a certain substance there is obtained from it a 86 ELEMENTARY PRINCIPLES OF CHEMISTRY substance which, when compared with the original in com- position, shows the loss of one atom of hydrogen and one of oxygen and the gain of one atom of chlorine that is, one atom of chlorine has been substituted for one of hydro- gen and one of oxygen, but no more than these two can be thus substituted. The theoretic interpretation of this is that one atom of hydrogen and one of oxygen are held in a peculiar manner not shared by the other atoms of hydrogen and oxygen. Furthermore, the derived sub- stance, C 2 H 3 OC1, does not show the property of substi- tuting a metal for the hydrogen ; therefore it is assumed that the hydrogen atom, thus associated with the oxygen atom in leaving the molecule, is the same that in the origi- nal substance was replaceable by a metal. This, the second step in solving the molecular structure, is expressed by writing the formula HO(C 2 H 3 0). 181/2 It is next shown by experiment that acetic acid can be made synthetically, by a series of changes not necessary to detail, from a substance whose molecule contains but one carbon atom, three hydrogen atoms, and the group OH. Its structure is shown by the formula CH 3 OH. The change of this into acetic acid involves the addition to the mole- cule, CH 3 OH, of a carbon atom from a source outside of itself, and some subsequent intermediate modifications ; but it is assumed that the atomic group CH 3 passes, itself unmodified, from the parent molecule, CH 3 OH, into the product, HO(C 2 H 3 0), and continues to exist in the latter. This resolves the atomic group C 2 H 3 into the groups CO 181/3 and CH 3 . Thus the original molecule containing two car- bon, four hydrogen, and two oxygen atoms (C 2 H 4 2 ) has been resolved into atomic groups as expressed by the sym- bols HO, and CO, and CH 3 . This is shown in the struc- tural or constitutional formula CH 3 -CO*OH. 182 By observations and assumptions, such as these just described, the structures of hundreds of substances have been determined, many of them very complicated. Not THE ATOMIC THEORY 87 only is this true, but also that many natural substances, and even many never found in nature, have been made artificially by building them up from simpler substances as indicated by the atomic groups in their assumed structure. It is doubtful if any branch of natural science can show more numerous instances of brilliant achievements in fact, realized under the guidance of theoretic conceptions, than can synthetic chemistry. 13. As to space relations of atoms, or stereo-isomerism. It is 183 assumed that the properties of substances may be influenced to a limited extent by the space relations of the atomic groups. The discovery made by Pasteur in 1848, that substances 184 might have not only the same percentage composition, the same molecular weight i. e., the same kind and number of atoms in the molecule but even the same atomic grouping, and still differ slightly in respect to certain properties, could not be explained by the then existing theories. The substance in which this phenomenon was first observed by Pasteur is tar tar ic acid, found in the grape and other fruits. The properties in which the slight difference is manifested are crystalline form, and behavior toward polarized light. It would be entirely impracticable fully to describe here either the phenomenon or the theory. Let it suffice to say that Pasteur observed at least two varieties of tartaric acid. One when in solution rotates to the right the plane vin which a ray of light is polarized ; the other rotates it to the left. The two varieties, when crystallized, show two kinds of crystals, one of which, as to arrangement of angles and faces, is like the image of the other as seen in a mirror. Yet, as to behavior in all their reactions, the two varieties are alike, and therefore, it must be assumed, have the same structure or atomic grouping. This phenomenon is called physical isomerism, or stereo-isomerism. Many other ex- amples of it have since been discovered. In 1874 Le Bel and van't Hoff, independently of each 185 other, announced a theory designed to account for the 88 ELEMENTARY PRINCIPLES OP CHEMISTRY phenomenon of physical isomerism. They both had the same fundamental idea. According to this, the carbon atom, which is capable of uniting with four hydrogen atoms, or four groups of atoms, as seen in methane, CH 4 , furnishes one condition for the phenomenon. It is assumed that these four atoms or groups are placed, with reference to the carbon atom, like the four apexes of a tetrahedron with reference to its center. Kow, it is evident that, if every one of four such atoms, or groups, is different from every other one, there might be in two molecules, otherwise alike, the same difference in the position of the four with reference to the carbon atom as there is in the four apexes of a tetrahedron with reference to its center, compared with the same as seen in the image of the tetra- hedron reflected in a mirror. A carbon atom thus linked with four different atoms, or groups, is called asymmetric, and its presence in the molecule is supposed by the the- ory to be the cause of stereo-isomerism. For example, the accepted structure of the molecule of lactic acid, the acid of sour milk, is shown by the formula H CH 3 C COOH, OH in which it is assumed that one carbon atom is linked with four different groups, viz., CH 3 , and H, and OH, and CO'OH. It is, therefore, asymmetric, and the substance should show the phenomenon of physical isomerism, as in fact it does. The theory has been admirably worked out, and finds strong support in many recognized facts. 186 Much more is included in this chapter than the atomic theory of Dalton contained, and more than this title accu- rately describes. But the title is classic, and serves its pur- pose probably as well as would any other single phrase. That these theories leave much unexplained, is surely true. HERMANN VON HELMHOLTZ B. Germany, 1821. D. 1894. (See No. 50, note.) THE ATOMIC THEORY 89 That they explain with entire satisfaction all that they undertake to explain, is not wisely claimed. That they may not in the future be greatly modified, and even es- sentially changed, it is not in accordance with the scientific spirit to assert. But that they have been and are useful as guide to investigation and as aid to understanding, is amply proved by the results achieved. CHAPTEE VIII RELATION BETWEEN THE PROPERTIES OF THE ELEMENTS IN GENERAL AND THEIR COMBINING WEIGHTS 187 IT is the purpose in this chapter to present a somewhat detailed description of the first twenty-five elements, taken in the natural order that is, the order of their increasing combining weights, and in connection with each one, for convenience of arrangement, to give attention to some of its important compounds. This is the subject-matter often designated as Descriptive Chemistry, and includes more, perhaps, than is implied, strictly speaking, in the title of the chapter ; nevertheless, the central idea of the whole will be found in the relation to which the title refers. A. Or THE ELEMENTS, COLLECTIVELY 188 All substances, so far as known, are made up of com- paratively few elementary forms of matter, which resist all attempts to reduce them to simpler forms. These, at present, number seventy-four, with possibly one or two additions. 189 Their distribution is by no means uniform. Either free or as constituents, all of them are found in the solid mass of the earth ; some thirty, in the sea ; five, in the atmos- phere, namely, hydrogen, carbon, nitrogen, oxygen, and argon (Roscoe) ; and, according to the revelations of the spectroscope, at least twenty-two and possibly thirty-eight have been identified in the sun (Eamsay). 90 DESCRIPTION OF ELEMENTS AND COMPOUNDS 91 In quantity, too, they differ greatly, so far as known in this fragment of the universe, the earth. Oxygen is the most abundant of all. It constitutes nearly one quarter (23 per cent) of the atmosphere, nearly nine tenths (88.9 per cent) of all the water of the globe, and nearly one half (Eoscoe) of the solid portion of the earth's crust, the other half being made up in the main of only seven other elements. The following is the estimated composition of the greater portion of the earth's crust, not including water (Eoscoe) : 190 Oxygen, 44.0 to 48.7 per cent Silicon, 22.8 to 36.2 " " Aluminium, 9.9 to 6.1 " " Iron, 9.6 to 2.4 " " Calcium, 6.6 to 0.9 per cent Magnesium, 2.7 to 0.1 " " Sodium, 2.4 to 2.5 " " Potassium, 1.7 to 3.1 " It is probable that the interior mass of the earth consists largely of sulphides (Eamsay). Adding to the eight of the foregoing list the following six, namely, hydrogen, carbon, nitrogen, phosphorus, sul- phur, and chlorine, we have those which constitute the greater part of matter as known to us. Twenty-three more might be named, which, with those already mentioned, thirty-seven altogether, would include about all with which we come in contact in everyday life. Of the remainder, some are found somewhat commonly, although in small quantities ; others are so rare as to be only scientific curi- osities ; and of still others it may be said that their very existence as elemental forms is open to question. The five most closely associated with the living organism, 191 whether plant or animal, are carbon, oxygen, nitrogen, hydrogen, and sulphur. Protoplasm, the fundamental form of living substance, contains these, and in proportion probably somewhat as follows : % Carbon, 51 to 55 per cent Oxygen, 20 to 24 " " Nitrogen, 15 to 17 " " Hydrogen, 6.5 to 7.5 " " Sulphur, 0.3 to 2 " " 92 ELEMENTARY PRINCIPLES OF CHEMISTRY Besides these, phosphorus, potassium, calcium, magne- sium, and iron are reckoned as essential to plant life. All of the foregoing are found in the normal human body, and in addition chlorine, fluorine, silicon, lithium, sodium, and manganese (Martin), If any one more than others could be regarded as the basis of living things, it is carbon, by reason both of its peculiar nature and its pre- dominating quantity. 192 Only comparatively few of the elements are found free, or uncombined, and abundant in nature, at least on the earth. Such are oxygen, nitrogen, carbon (as coal), and sulphur. Less abundant, yet familiar either as natural or artificial products, are some of the metallic elements, namely, magnesium, aluminium, iron, nickel, cobalt, zinc, copper, tin, lead, mercury, silver, gold, and platinum. 193 Of the whole list of elements the following are generally classed as distinctly non-metallic : boron, carbon, silicon, nitrogen, phosphorus, oxygen, sulphur, selenium, fluorine, chlorine, bromine, iodine, and the newly discovered sub- stances helium and argon, fourteen in all. The following are classed sometimes as metallic, sometimes as non-metallic : hydrogen, titanium, zirconium, vanadium, arsenic, anti- mony, and tellurium. The rest are metallic. 194 As to physical properties, they vary through an enor- mous range. At ordinary temperature, seven are gaseous, namely, hydrogen, nitrogen, oxygen, fluorine, chlorine, he- lium, and argon. Two are liquid bromine, and mercury. 195 The others are solid. In boiling point they vary from that of hydrogen, 238, to that of carbon, which volatilizes only at the highest temperature of the electric furnace, that is, 3,500 or more. 196 In density, the extremes are hydrogen, one cubic centi- meter of which weighs 0.0000899 of a gram, and osmium, of which the same volume weighs 22.48 grams; that is, osmium weighs 250,000 times as much as hydrogen, volume for volume. DESCRIPTION OF ELEMENTS AND COMPOUNDS 93 In combining weight they range from 1 for hydrogen to 197 238 for uranium. In chemical activity, also, they vary greatly. Thus fluo- 198 rine is so reactive, attacking so energetically everything with which it comes in contact, that, although its existence was recognized, it could not be isolated until Moissan ac- complished the feat in 1886. On the other hand, argon, although present in the atmosphere to the extent of one per cent, escaped detection until 1894, perhaps because of its great inactivity, as all of the many attempts to make it enter into combination have been so far (1899) without established success. Oxygen is the most universally react- ive, forming compounds with all the other elements, except fluorine, argon, and helium. The heat disturbance per gram of element in combining 199 with oxygen ranges from 34,200 calories for hydrogen in the formation of water, to 1540 calories for nitrogen in the formation of nitric oxide; or if the comparison is made per combining weight, which is probably more suitable, the values are 143,900 calories for 24 grams of magnesium, and 21,600 calories for 14 grams of nitrogen. 1. HYDROGEN Symbol H. Comb, wt. 1 History. Hydrogen was studied and for the first time identified by 200 Cavendish in 1766, although Paracelsus in the sixteenth century noted the production of an inflammable gas by the action of acids on metals. In 1781 Cavendish showed that it is a constituent of water. Its name signifies water-producer, and was given it by Lavoisier. Natural occurrence. It is found uncombined, but only in 201 small quantities, in volcanic gases, in the gas from oil wells, and sometimes in meteorites. On the other hand, as a constit- uent, hydrogen is both very abundant and widely distributed. It forms one ninth of all water, and is contained in organic matter of both plant and animal origin, and also in all acids. 94 ELEMENTARY PRINCIPLES OF CHEMISTRY 202 Preparation. 1. By the passage of the electric current through slightly acidulated water (electrolysis), by which hydrogen and oxygen are set free. FIG. 2. The electrolysis of water. Showing how the current from the electric battery is passed through acidulated water, and the gases col- lected in used for firearms as early as 1346. It is a mixture of potassium nitrate, charcoal, and sulphur. The ingredients are carefully purified, finely ground, and very thoroughly mixed. The mixture, slightly damp, is put under heavy pressure from 300 to 450 pounds to the square inch in order to increase its density. The " press-cake " thus obtained is broken into small grains " granulation " the process is called and sifted. The grains are smoothed and glazed by rolling in wooden drums, sometimes with the addition of a little graphite. The product is again sifted, and freed of moisture by drying at a low temperature. A black powder, commonly used for military purposes, contains 75 per cent of potassium nitrate, 15 per cent of charcoal, and 10 per cent of sulphur. Some idea of the chemical change in the explosion is given 55 7 by the incomplete equation : 2KN0 3 + 3C + S = 3CO + 2N + K 2 S + . 206 ELEMENTARY PRINCIPLES OF CHEMISTRY But it is actually, at least when taking place in a confined Space, much more complicated. However, it is clear that the carbon and the sulphur may be converted into gaseous oxides by the oxygen of the nitrate, and thus combustion take place without the intervention of air to any extent. The large volume of the resulting gases and the heat liberated by the reaction cause the energy of the explosion. The gaseous products include : Carbon dioxide, C0 2 ; Nitrogen, N ; Carbon monoxide, CO ; Hydrogen sulphide, H 2 S ; Hydrogen, H ; Methane, CH 4 . The hydrogen may come from the charcoal or from the moisture pres- ent. Of these gases, the first three make about 90 per cent of the total, while somewhat more than 90 per cent of the total solid products are made up of Potassium carbonate, K a C0 8 ; Potassium sulphate, KaSO* ; Potassium persulphide, K 2 S 2 . 558 The reaction is expressed with closer approximation by the equation, 8KN0 3 + 9C + 3S = 2K 2 C0 3 + K 2 S0 4 -f- K 2 S 2 + 7C0 2 -f 8N. From 1 gram of powder the solid products weigh from 0.55 to 0.58 of a gram, and the gaseous from 0.45 to 0.42 of a gram, while the vol- ume of the gases reckoned at is from 200 to 300 cubic centimeters, and the temperature is estimated at 2,100 or 2,200. 559 It is evident that the rending effect of an explosion must depend upon the rapidity of the reaction, for if the latter is slow the gases are slowly liberated, and the heat is gradually dissipated, so that the tem- perature does not rise as high ; moreover, the pressure of expansion, which when suddenly applied may be irresistible, when slowly applied may be resisted. Now, if the reaction is one of oxidation the rapidity must be influenced by the closeness of contact, or intimacy of mixture, of the combustible with the supporter of combustion. It has been seen in the experiments (see Exp. SSI/,) that a piece of ignited char- coal burns actively, but not explosively, when brought in contact with melted nitrate. It is quite different if the two substances are finely powdered and then intimately mixed, as in gunpowder, each small grain of which contains its proportion of the three ingredients. Again, if the grains are made larger and therefore do not lie so closely in con- DESCRIPTION OF ELEMENTS AND COMPOUNDS 207 tact, the rate of combustion through the whole mass, when the ignition is started at only one point, must be retarded. For this reason the powder used in heavy guns, known as " pebble " and " prismatic " powder, is made into grains which measure from five eighths to one and three quarters inches, so that the full pressure of the explosion is not reached at once. Were it otherwise, the gun might yield before the ball was set in motion. Again, the velocity of combustion through a mass of gunpowder is 560 influenced by the pressure of surrounding gases. In a vacuum it does not explode at all, but burns slowly. In open air the ignition travels through the mass at the rate of four feet per second ; in a heavy gun at about thirteen feet per second. So it happens sometimes, in using large-grained powder, that the charge is ignited and the ball leaves the gun before combustion reaches the last of the powder, and this is thrown out unburned. Likewise in exploding powder under water, if the containing case is not strong enough to hold the gases until the ignition is complete, a part of the charge may remain unexploded (Walke). By varying the quality of the charcoal, a " brown " or " cocoa " powder is made which is so slow in burning that a large grain may be ignited while held in the hand and dropped before th'e fire reaches the fingers. Still again, the velocity of the combustion is influenced by the man- 561 ner of firing. In the foregoing statements it has been implied that the chemical reaction which produces the explosion is started by a suffi- cient rise of temperature in only a small portion of the powder, such as is practically caused by a spark of fire or a wire made hot by the electric current; and this ignition, producing heat and flame, spreads with more or less rapidity through the whole. But there is another mode of 562 firing, known as detonation. In this the reaction is brought about by the shock of another explosion and can not be due, at least not in all cases, to rise of temperature. A very common detonating material is mercury fulminate, the nature of which need not be explained here. Its application is seen in the ordinary percussion cap in which it is exploded by a blow, although when this is used in connection with powder the latter is probably fired by the small flame or spark from the cap, and so, strictly speaking, is not detonated. However, another explosive namely, gun cotton even when wet may be exploded by the explosion of a small quantity of fulminate in contact with it. In this case it can not be rise of temperature which explodes the gun cotton. This shock of detonation is brought to bear upon the whole mass of the explosive within the reach of its influence instantaneously, or at least with a rapidity very much greater than the velocity of ignition. 208 ELEMENTARY PRINCIPLES OF CHEMISTRY It is as if a violent blow were given to every minutest particle of the explosive at practically one instant of time, and therefore the explosion is made more violent. Gunpowder is not readily detonated ; indeed, this method of explosion is not desirable for firing projectiles, and it is somewhat uncertain whether this material can be really detonated in any conditions. There is a class of explosives of altogether different chemical character which are most readily detonated. Before consider- ing these, some modification in the constituents of gunpowder may be mentioned. 563 The nitrate of sodium, because of its cheapness, is substituted for the potassium salt in mining powder, especially when it is to be used in hot and dry countries. It was largely employed in building the Suez Canal, but its hygroscopic quality hinders its general use. The same objection applies to ammonium nitrate, although it has the advan- tage of producing a larger volume of gases and less solids, therefore less smoke. 564 Many attempts have been made to utilize potassium chlorate as oxi- dizing material. This decomposes at about 350 into potassium chlo- ride and oxygen, with liberation of heat. A chlorate powder produces higher temperature and larger volume of gas, and simpler, less disso- ciable products ; the reaction begins at lower temperature and spreads with greater rapidity, and therefore causes more abrupt and shattering effect than with nitrate powders. But the former are much more sensi- tive, and are subject even to (so-called) spontaneous explosion ; therefore much greater danger attends their preparation, storage, and use. Be- sides, the products of decomposition are more corrosive of the gun and more injurious to the person than with the nitrate powders. The dis- advantages overbalance the advantages, and have prevented the prac- tical success of powder of this class for firing projectiles. But mix- tures containing chlorate are used for charging explosive shells and bullets, and for fuses which are to be ignited by friction, percussion, or .contact; in them, sugar and starch are sometimes used as the com- bustible material. 565 Liquid mixtures. Some attempts have been made to use liquid mix- tures of combustible and oxidizer, in which evidently the contact of the two would be more intimate than in the mechanical mixture of powders. One such mixture consists of nitric acid as oxidizer and sol- vent, and certain hydrocarbons as combustible. They seem to have no practical success. 566 Gaseous mixtures must bring the combustible and oxidizer into inti- macy of the same order, as in liquids, reaching to the minutest par- ticles of the substances that is, to the molecules themselves. But, DESCRIPTION OF ELEMENTS AND COMPOUNDS 209 from the very nature of the gaseous condition, it would seem that reac- tion by shock could not be brought about as in liquids and solids, and that the spread of the combustion must depend on temperature or in- flammability. In combustion of this kind the volume of the products is not necessarily larger than that of the explosive mixture ; indeed, it may be less, as when three volumes of hydrogen and oxygen become two of water. The explosive power then must be due to the rise of temperature in the gases, which is caused by the heat of reaction. Gaseous mixtures are not applied to the ordinary purposes of explo- sives, but they are often the cause of serious accidents, and they find a limited use as motive power in gas engines. Another type of explosive mixture has likewise been the cause of 567 most serious accidents namely, a combustible solid in the form of very fine dust scattered through the air in a more or less confined space. It will readily be understood that minute particles of carbon might be suspended in air so close to each other that ignition, started at one point, would flash quickly to neighboring points exactly in effect as if it were a gaseous mixture. In such way most disastrous explosions have occurred in flour mills and starch factories, where the extremely fine dust of combustible starch, swept into the air by some accidental cause, and brought in contact with flame, has furnished the necessary conditions. The dust in coal mines also has caused explosion. Explosive compounds. Finally, there is still another and most inter- 568 esting type of explosive, in which the closest kind of contact is brought about. In these the combustible and the oxidizer are proximate con- stituents of one individual substance. In theoretical terms, they are parts of the same molecule instead of being in neighboring molecules. Therefore the explosive is a single substance and not a mixture. How this can be, is understood in recalling the fact that of potassium nitrate, KaO'N 3 5 , it is mainly the nitrogen oxide which supplies oxygen to the carbon in gunpowder. Now, if nitric acid could be combined in a salt with a combustible base for example, one containing carbon and hydrogen the condition of combustible and supporter of combustion would be realized in one compound, and it would only be necessary to break down the compound in order that the elements might recombine into carbon dioxide, water, and nitrogen. Such a base is found in common glycerine, which is only one of a numerous class of organic bases, called alcohols. The composition of glycerine is seen in the formula, C 3 H 6 (OH)g, and its combination with nitric acid, after analogy with the formation of potassium nitrate, is shown by the equations : C,H.(OH) + 3HNO, = C 3 HB(N0 3 )3 + 3H a O. KOH + HN0 3 = K(N0 3 )+ H 8 0. 210 ELEMENTARY PRINCIPLES OF CHEMISTRY 569 This glycerine trinitrate is more commonly known under the less cor- rect name of nitroglycerine. The following equation shows that no outside oxygen is needed for the complete combustion : 2C 3 H 6 (N0 3 ) 3 ) = 6CO a + 5H 2 + 6N + 0. Nitroglycerine was discovered in 1847, but was practically unused until about 1863. Its successful application was due to a Swede, Nobel by name, who also discovered the method of firing by detonation. Nitro- glycerine is made by adding glycerine, as a thin stream or spray, very slowly to a cooled mixture of sulphuric and nitric acids in the most concentrated form obtainable. It is important that all these substances be very pure and anhydrous. The temperature of the mixture must not rise above 30 ; if it does, there is great danger. When the reaction is over, the nitroglycerine rises to the top and is drawn off, or the mixture is run into water in which the product sinks to the bottom as an insolu- ble oil. It must then be very thoroughly washed until free of acid, as the presence of this promotes decomposition and dangerous instability. It is sometimes strained through felt filters for additional purification. 570 Nitroglycerine is an oily, colorless, or nearly colorless liquid, heavier than water, with which it does not mix. It is somewhat volatile at 50 and freezes at about 8, although the freezing point varies in different samples. It is poisonous, and is used as a powerful remedy in medicine. In the open air, small quantities may be burned without explosion, but when heated to about 180, it explodes. It may also be exploded by a shock, either of a blow or by detonation. It is less sen- sitive in the frozen condition than in the liquid, and it is not commonly used in the latter form, except in order to " torpedo " gas and oil wells. 571 The volume of gas (counting the water as gaseous) from the explo- sion of one gram of nitroglycerine is about 714 cubic centimeters reck- oned at 0, and the temperature is about 3,000. It is reckoned as from four to six times more effective than powder, at least in blasting rock. The velocity of detonation in the liquid is 5,300 feet per second. Owing to its properties, nitroglycerine does not need to be confined in order to be effective. Exploded on the surface of a rock, it may shat- ter the rock, for the action is so quick that the atmosphere is practi- cally unyielding during the short interval. 572 In order to reduce the danger in handling liquid nitroglycerine, it is absorbed in some solid which in one class of explosives is inert, in an- other class the absorbent is itself explosive, or at least combustible. Of the first class, and the most common of all, is dynamite, which was in- vented by Nobel. The absorbent is a kind of clay made up largely of the silicious remains of minute organisms. This takes u-p about three DESCRIPTION OF ELEMENTS AND COMPOUNDS 211 times its weight of nitroglycerine and forms a plastic mixture. The velocity of detonation in dynamite is 20,000 feet per second, four times that in the liquid, and its effective intensity is greater, but in smaller ratio ; these facts are difficult of explanation. As absorbents, mag- nesia, powdered mica, sawdust, and charcoal are used to some extent. Gun cotton is an explosive of the same chemical type as nitro- 573 glycerine. It is a nitrate of cellulose ; cellulose is the alcoholic, organic base which constitutes the fundamental substance of the plant struc- ture, and it is seen in nearly pure condition as cotton fiber and filter paper. The nitrate is made from thoroughly cleaned cotton by dip- ping it into the mixture of concentrated sulphuric and nitric acids. In this instance, also, it is very important that the acids be completely washed out, lest they provoke dangerous decomposition ; therefore the cotton, which is hardly changed in appearance by nitrating, is reduced to pulp under water, then washed and pressed into compact forms as desired. When designed for military purposes, it is allowed to retain from 16 to 30 per cent of water, which renders it much safer, and does not interfere with its explosion by detonation. In this condition it is stored, for it is not sensitive to friction or percussion, nor to fire until the water is dried out of it. Even when dry it is not easily exploded by percussion, and it burns quietly if unconfined. When properly pre- pared, it is reckoned as " the safest explosive known." (Walke.) The composition of gun cotton, reduced to the simplest terms, is 574 expressed by the formula C 6 H702(N0 8 )3. The following equation shows that it does not contain enough oxygen for complete combustion : 2C 8 H 7 2 (N0 3 ) 3 + 90 = 12C0 2 + 7H 2 + 6N. In the gaseous products of its explosion (there are no solids, hence no smoke) are found carbon monoxide and hydrogen, besides those in the preceding equation. The total volume of gas from one gram of gun cotton is given by one observer as 859 cubic centimeters, reckoned at 0. This is more gas than either gunpowder or nitroglycerine gives, and the temperature produced is also higher than that of gunpowder. The velocity of detonation is greater in wet gun cotton than in dry, reaching about 20,000 feet per second, nearly four times the velocity in nitroglycerine and about the same as that in dynamite. Gun cotton is used largely as a high explosive for military purposes, 575 in torpedoes, and in submarine mines. It is also used as a constituent in many other explosives. In some, additional oxygen is supplied by mixing with potassium or other nitrate. Blasting gelatin is a mix- ture of mono- and di-nitrocellulose and nitroglycerine, which contains some excess of oxygen. (See No. 569.) Nitrocellulose is also the chief 15 212 ELEMENTARY PRINCIPLES OF CHEMISTRY 57(5 constituent of the smokeless powders which are displacing the black and the brown gunpowder for military and other uses. That which is prepared for the United States navy contains nitrocellulose mixed with barium nitrate, potassium nitrate, and calcium carbonate. Cordite, the powder adopted by Great Britain, contains nitroglycerine, nitro- cellulose, and vaseline (Deering-Thorpe, F. H. Thorp, Walke). 17. CALCIUM Ca. 39.7 577 History. The preparation of lime, CaO, by the so-called burning of limestone, CaC0 3 , and its use in making mortar, are of great an- tiquity. In 1756 Black showed the chemical relation between lime and limestone, and in 1808 Davy obtained the element as a metallic powder by the electrolysis of the chloride. Matthiessen, in 1856, also by elec- trolysis, succeeded in producing it as a coherent metal, and Moissan re- investigated it in 1898. 578 Natural occurrence. It is never found free. Its most abundant compound is the carbonate, which is known in its various conditions as limestone, chalk, marble, coral, calc- spar, calcite, etc. The sulphate, CaS0 4 , is also abundant under the names of anhydrite and gypsum. The phosphate, borate, silicate, and fluoride are native. Calcium salts are present in natural waters, and are essential to the plant organism, accumulating in the leaves, and also essential to the animal, occurring particularly in shells, bones, and teeth. The element is present in the sun, and has been found in meteorites. 579. Preparation. Moissan prepared it (1898) by electro- lyzing fused calcium iodide, CaI 2 , at a dull-red heat, also by heating to dull red a mixture of the anhydrous iodide and sodium in a closed iron crucible. The reaction is : CaI 2 + 2Na = Ca The iodide and the excess of sodium are dissolved out by anhydrous alcohol, which has been freed of dissolved air, and the calcium is left as a crystalline powder. DESCRIPTION OF ELEMENTS AND COMPOUNDS 213 Properties (Moissan). Calcium is a silver-white, crystal- 680 lizable metal, soft enough to be cut with a knife. Its spe- cific gravity is 1.85, and it melts at 760. Heated in oxy- gen to 300, it burns brilliantly, with great evolution of heat, forming the oxide CaO. Also in nitrogen it burns, forming the crystallizable nitride Ca 3 N 2 . Both oxide and nitride are formed when it burns in air. Heated in hydro- gen, a crystalline hydride, CaH 2 , is formed ; and in chlorine, the chloride, CaCl 2 . It combines also directly with carbon (CaC 2 ), with silicon, with sulphur, and with phosphorus (Ca 3 P 2 ). A compound with boron, and a second oxide, Ca0 2 , are known. Calcium decomposes water on contact at ordinary temperature, liberating hydrogen and forming the oxide CaO, which combines with water and acts as base in forming salts. The metal reduces the oxides of lithium, sodium, and potassium, also carbon dioxide, when heated with them, but not the oxide of magnesium. The hydride, carbide, nitride, and phosphide decompose water on contact by reactions, which are thus expressed : CaII 2 + 2H 2 = CaOH 2 + 4H. CaC 2 + 2H 2 = CaOH 2 + C 2 H 2 . 6H0 = 3CaOH0 + 2NH 8 . Lime, calcium monoxide, CaO, does not occur native. It 581 is made on a large scale by heating the carbonate, CaC0 3 ; the decomposition begins at about 400. This is called technically, but erroneously, "burning," and the product " burnt " lime, or " quicklime." Lime is a white and, ordi- narily, amorphous solid. Exposed to prolonged heating by the oxyhydrogen flame, it slowly crystallizes at the surface, but by the heat of the electric arc it fuses and volatilizes, forming colorless, brilliant crystals. If impurities e. g., alumina are present, fusion takes place more readily, since compounds with the lime are formed which are more fusi- ble. The amorphous lime which is made from the carbon- ate is hard and quite porous ; by exposure to air it absorbs 214 ELEMENTARY PRINCIPLES OF CHEMISTRY water and carbon dioxide slowly, and falls into a soft pow- der, which is a mixture of hydroxide and carbonate. This makes quicklime a useful drying substance, especially for gases. On the other hand, if lime is brought in direct con- tact with a suitable quantity of water, combination takes place quickly with liberation of much heat ; the hydroxide, Ca(OH) 2 , is formed, and appears as a soft dry powder, if too much water is not present. This operation is called " slaking," and the product " slaked " lime. The heat of hydration may be sufficient to cause fire if combustible matter is present. The crystallized oxide, on contact with air or with water, changes very much more slowly than the amorphous, and apparently it does not dissolve until the combination with water has taken place. 582 The hydroxide, Ca(OH) 2 , usually an amorphous powder, although crystallizable, is slightly soluble in water, 100 grams of the latter dissolving about 0.18 of a gram of the hy- droxide at 20, but only one half as much at 100. The solubility, therefore, decreases with rise in temperature, which is exceptional. The solution has alkaline reaction and is often called "limewater." When more hydroxide is present than the water can dissolve, the mixture is called " milk of lime." The hydroxide loses its water at a red heat, the decomposition beginning even at 100 (Kamsay). 583 Lime finds a great variety of uses, for instances : in mortar and cements ; in making glass, bleaching powder, and soda (by the Leblanc process) ; in purifying coal gas and sugar ; in preparing many chemicals ; in bleaching and dyeing cotton fabrics; in tanning leather; in obtaining metals from their ores ; and as a disinfectant. 584 The monoxide and the hydroxide act as base and form a series of well known salts. The chloride, CaCl 2 , when anhydrous is very deliquescent, and dissolves abundantly in water with evolution of heat. The crystallized chloride, CaCl 2 -6H 2 0, dissolves with absorption of heat, and, mixed with ice or snow, may reduce the temperature to 40 ; this DESCRIPTION OF ELEMENTS AND COMPOUNDS 215 makes a convenient freezing mixture. Bleaching powder, and the carbonate, also the phosphate, have been sufficiently considered in other connections. The sulphide,. CaS, in the impure form commonly made, is phosphorescent and used somewhat in making the luminous paints. The sulphate, CaS0 4 , known as " plaster of Paris," is used in cements. 17a. Mortar and other Cements Preparation of lime. The primitive limekiln or furnace was a pit 585 sunken in the ground of a hillside ; into it was piled the limestone in rather large lumps, leaving an opening and cavity at the bottom for the fuel. The hot gases and the carbon dioxide which is liberated are then free to pass upward through the loosely packed material and escape into the atmosphere. It is important that the carbon dioxide be thus swept away as produced, otherwise the carbonate may be reformed. Permanently constructed furnaces are now largely used. Excessive heating must be avoided, lest it cause the beginning of fusion , which gives a product that does not easily slake. This is more likely to happen with impure limestone. For other reasons also, approximate purity is desirable in many of the uses to which lime is put. Mortar is a mixture of lime and sand which dries in the atmosphere 58(> and hardens ; it is used as a cement, as a covering for walls, and in many other ways. In its preparation, the quicklime is first slaked and mixed with water to a thin paste. If such a paste is allowed to dry, it shrinks considerably ; therefore it is mixed with sand, and in this con- dition it is applied to brick or stone, so that it fills the separating spaces completely, and hardens to a compact mass without shrinkage. Sand, the grains of which are sharp and not rounded by friction, is preferred, as experience has shown that it makes a better mortar. After several days of exposure to air the mortar " sets " that is, partly hardens. This is due to drying. After this change, a very slow absorp- tion of carbon dioxide takes place, with the formation of carbonate, and this produces the second stage of hardening. The sand serves not only to increase the bulk of the lime, but also its porosity, and thus facilitates the absorption of carbon dioxide. If the drying is too quick or too complete, the hardening is interfered with, so that moisture plays an important part in the change ; just how this is, and how the carbon dioxide penetrates to the interior, are not entirely clear. It is thought that the sand is not chemically changed in hardening, although in samples of mortar one hundred years old, or older, the silica, from 2 216 ELEMENTARY PRINCIPLES OP CHEMISTRY to 6 per cent of it, has been found in combination. The mortar in the Pyramid of Cheops, although older than 2000 B. c., is practically of the same composition as the modern mixture ; and other mortars of Phoe- nician, Greek, and Roman masonry have been examined with similar result. In some instances the lime is found converted completely into carbonate, in others only partly so. Hydraulic lime. If limestone contains more than about 10 per cent of clay (aluminium silicate) when it is burned, it yields a lime which slakes less readily and with less heat, but which sets or hardens when in contact with water or even under water ; hence it is called hydraulic. The hardening in this case is due to reaction between water and the anhydrous silicates and is independent of carbon dioxide. Such lime is of value in making hydraulic mortar or cement which is to be used in constructions with which water comes in contact. Certain natural anhy- drous silicates of volcanic origin, when finely ground (without burn- ing) and mixed with ordinary lime, constitute another variety of hy- draulic cement; and blast-furnace slag, if cooled quickly from the melted condition, is similarly used with lime. Still a third variety is made by burning at a very high temperature an artificial mixture of pulverized calcium carbonate and clay. Of this variety is " Portland cement," which is made and also imported in large quantity in this country. The materials are ground and intimately mixed, then burned, and finally ground again. Much of the quality depends upon the proper heating. The setting and hardening are undoubtedly due to chemical reaction, but its exact nature is not easily explained. It is supposed that silicate and aluminate of calcium are formed by the heating, and that these combine with water and crystallize more or less, finally becoming as hard as natural stone. Plaster of Paris is another variety of cement which finds many uses. This is made by heating to 120 or 130 the mineral, gypsum, which is a hydrated calcium sulphate, CaS0 4 -2H 2 0. Only about three fourths of the water is thus expelled, so that plaster of Paris is chemically (CaS0 4 ) 2 -H 2 0. When this powder is mixed with water to a creamy paste, heat is liberated, the mixture swells somewhat, and quickly hard- ens to a white porous mass of smooth surface. If the gypsum is heated to 200, all the water is driven off and the plaster does not set. The chemical changes of the hardening are explained as follows : The com- pound (CaS0 4 ) 2 -H 2 is somewhat soluble, and in solution slowly com- bines with water, thus : (CaS0 4 )iH,0 + 3H,0 = 2(CaS0 4 -2H 2 0). The latter hydrate, being less soluble than the first, separates in minute interlacing crystals. Plaster of Paris is used as a wall finish in the inte- DESCRIPTION OF ELEMENTS AND COMPOUNDS 217 rior of buildings, also for making casts and reproductions for which the property of expanding in setting makes it especially applicable (Hartley-Thorpe and F. H. Thorp). REVIEW PROBLEMS 1. Assume that magnesium is converted into its iodide (see Exp. 41/b) by direct action, how much iodine is theoretically needed, and how much should the magnesium iodide weigh f 2. Suppose that one gram of magnesium is to be converted into magnesium pyrophosphate and as such weighed ; suggest the practicable operations for so doing, writing them out in the form of directions for the experiment, based on your knowledge of the properties of the sub- stances involved and your experience in experimentation (see Exps. 446A and 494). 3. Calculate what should be the weight of the pyrophosphate thus obtained. 4. Suppose that dry magnesium sulphate is to be the starting point for the same result as in 2, in what respect would you modify your directions'? What quantity of the sulphate should yield the same quantity of the pyrophosphate that one gram of magnesium yields ? 5. Does aluminium easily burn ? Suggest practicable steps by which one gram of aluminium may be converted into the oxide by converting first into the hydroxide. How much should the oxide thus obtained weigh I 6. What is the ratio between those masses of magnesium and aluminium which exactly react with equal quantities of hydrochloric acid f Also with equal quantities of oxygen f 7. How much iodine should be needed exactly to convert sulphur- ous acid into sulphuric acid, making enough of the latter to neutralize exactly 39.8 grams of pure sodium hydroxide ? 8. Sulphuric acid may be completely precipitated from solution by adding sufficient barium chloride solution ; barium chloride in crystal- lized condition has the formula BaCl 2 -2H 2 : calculate how much of the latter would be necessary to precipitate exactly the sulphuric acid of 1 gram of sodium sulphate crystallized (Na 2 S0 4 -10H 2 0)? 9. How much crystallized silver nitrate (AgN0 3 ) is necessary to precipitate exactly the chlorine of 1 gram of dry sodium chloride f 10. What are the relative masses of chlorine and of anhydrous nitric acid (HN0 3 ) which are necessary to convert equal quantities of ferrous sulphate, FeO-S0 3 , into ferric sulphate, Fe a 3 -3S0 3 ? (See Nos. 532 and 317, Part I.) 11. As between sodium nitrate and potassium nitrate at the same 218 ELEMENTARY PRINCIPLES OF CHEMISTRY price per kilogram, which is the more economical material with which to effect oxidation, other things being equal ? What is the ratio of economy f 12. As between magnesia and lime at the same price per pound, which is the more economical material for liberating ammonia from ammonium chloride 1 What is the ratio of economy ? G. GENEKAL SUBVEY 589 The general survey made of the first nine elements may now be extended to include the seventeen. A study of the fuller data of Table VII, No. 590, brings out in a striking manner the peculiar progression in properties at first only suggested. Thus as to melting point, beginning with so- dium, and again with potassium, one sees a progression like that which begins with lithium. The same may be said as to the combining volume, as to the heat of formation, the valence, and the basic and acidic function of the oxides. Especially with regard to valence and the basic function is clearly seen the tendency to recurring similarities at regu- lar intervals or periods. Beginning with the alkali metal lithium, valence increases by units to a maximum of four, then decreases to one, and starting again at one with the alkali metal sodium, it passes again through the same values to reach one in the alkali metal, potassium. In the same interval, the basic property diminishes, disappears, and re- appears. This is the germ of what is known as periodicity. Before giving a more detailed study to this important topic, it is well to consider some facts descriptive of five addi- tional elements which may be presented collectively. DESCRIPTION OF ELEMENTS AND COMPOUNDS 219 1. Ml ombin volum aa 590 It Oi r-< l> 00 JO rHt->i-l05Tt C^CO | TH 1 1 I+ + " CO CO ^JH" 50" r-j GO OS 05 GO CO 00 rH SSC, Manganese Mn 54 57 55.0 67 Mercury . Hz 198.5 200.0 THE ELEMENTS IN ALPHABETICAL ORDER 241 TABLE X. The Elements in Alphabetical Order (continued] No. NAME. Symbol. Combining weight. H = l. Combining weight. O=16. 33 Molybdenum Mo 95.3 96 54 Neodymiurn , Nd 142.5 (?) 143.6(?) 75 Neon 9.5 24 Nickel Ni 58.24 58.7 6 Nitrogen N 13.93 14 04 63 Osmium Os 189 6 191 15.88 16.00 41 Palladium Pd 105.6 106.4 13 Phosphorus . . P 30 8 31 65 ' Platinum Pt 193 3 194 8 16 Potassium K 38 82 39 11 58 Praseodymium Pr 139 4 (1} 140.5 (?) 40 Rhodium . ... Rh 102 2 103 33 Rubidium Rb 84 78 85 43 39 Ruthenium Ru 100.9 101.7 55 Samarium Sm 149. (?) 150 (?) 18 Scandium . .... Sc 43.8 44 1 31 Selenium Se 78.5 79 1 12 Silicon . Si 28 2 28 4 49 Silver Ag 107.11 107.92 9 Sodium Na 22.88 23 05 34 Strontium . . Sr 86.95 87 6 14 Sulphur S 31.83 32.07 61 Tantalum Ta 181.6 183.1 48 Tellurium Te 126.5 127.5 57 Terbium Tb 158 8 (?) 160 (?) 68 Thallium Tl 202.6 204.1 71 Thorium Th 230.8 232.6 59 Thulium Tu 169.4 (?) 170.7(?) 45 Tin Sn 118 1 119.0 19 Titanium Ti 47.8 48.1 69 Tungsten W 183. 184.4 79 Uranium u 237 8 239.6 90 Vanadium V 51 51.4 60 Yt 171.7 173. 35 Yttrium Y 88.3 89.0 97 Zinc Zn 64.9 65.4 36 Zirconium . Zr 89.7 90.4 242 ELEMENTARY PRINCIPLES OF CHEMISTRY 644 TABLE XI. The Elements in Natural Order No. NAME. Symbol. Equivalent weight. Fac- tor. Combining weight.* H = l. Combin- ing wt.* = 16. 1 Hydrogen . . . H 1. 1 1. 1.008 9 Lithium Li 6.97 1 6.97 7.03 3 4 5 6 Glucinum or beryllium Boron Carbon Nitrogen Gl B C N 4.5 3.62 2.9775 4. 43 2 3 4 3 9.0 10.86 11.91 13.93 9.1 10.95 12.00 14.04 7 Oxygen 7.94 2 15.88 16.00 8 Fluorine F 18.91 1 18.91 19.06 9 Sodium Na 22.88 1 22.88 23.05 10 11 12 Magnesium Aluminium .... Silicon Mg Al Si 12.05 8.967 7.05 2 3 4 24.10 26.9 28.2 24.28 27.1 28.4 13 14 Phosphorus .... Sulphur P S 10.267 15.91 3 2 30.8 31.83 31.0 32.07 15 Chlorine Cl 35.18 1 35.18 35 . 45 16 17 Potassium Calcium K Ca 38.82 19.9 1 2 38.82 39.7 39.11 40.0 18 Scandium Sc 14.6 3 43.8 44.1 19 Titanium Ti 11 95 4 47 8 48 1 90 Vanadium V 17.0 3 51.0 51 4 21 22 9,3 Chromium ... . Manganese Iron Cr Mn Fe 17.25 27.28 27 80 3 2 2 51.74 54.57 55 6 52.14 55.0 56 94 Nickel . . . Ni 29 12 2 58 24 58 7 9,5 Cobalt Co 29 3 2 58 6 59 96 Cu 31 56 2 63 12 63 60 97 Zinc Zn 32 45 o 64 65 4 98 Gallium Ga 23 13 3 69 4 70 29 30 Germanium .... Arsenic Ge As 17.975 24 8 4 3 71.9 74 4 72.5 75 31 39 Selenium Bromine Se Br 39.2 79 35 2 1 . 78.5 79 35 79.1 79 96 33 Rubidium Rb 84 78 1 84 78 85 43 34 35 Strontium Yttrium Sr Y 43.47 29.4 2 3 86.95 88.3 87.6 89.0 36 Zirconium Zr 22.43 4 89.7 90.4 * According to Clarke, Journal of the American Chemical Society, xxi, 213, February, 1899 ; and Richards, American Chemical Journal, xx, 543 (July, 1898). THE ELEMENTS IN NATURAL ORDER 243 TABLE XL The Elements in Natural Order (continued) No. NAME. Symbol. Equivalent weight. Fac- tor. Combining weight. Combin- ing wt. O-= 16. 37 Columbium or niobium Cb 31. 3 93 94 38 39 40 Molybdenum . . . Ruthenium Rhodium Mo Ru Rh 47.6 50.45 51.1 2 2 2 95.3 100.9 102.2 96.0 101.7 103 41 Palladium Pd 52 8 2 105 6 106 4 42 Silver 107 11 1 107 11 107 92 43 44 Cadmium Indium Cd In 55.7 37 67 2 3 111.4 113 112.2 113 9 45 Tin Sn 29 5 4 118 1 119 46 Antimony Sb 39 7 3 119 1 120 47 Iodine I 125 89 1 125 89 126 85 48 Tellurium Te 63.25 2 126.5 127.5 49 Cassium Cs 131.9 1 131.9 132.9 50 Barium Ba 68.20 2 136.39 137.4 51 59. Lanthanum .... Cerium La Ce 45.81 34.75 3 4 137.45 139.0 138.5 140.0 53 54 55 Praseodymium. . Neodymium Samarium Pr Nd Sm . 46.5 47.5 49.7 3 139.4(?) M49. (?) 140. 5(?) 143.6(1) 150. (?) 56 57 Gadolinium . . . Terbium . . . Gd Tb 51.87 52 93 3 3 155.6 (?) 158 8 (?) 156. 8(?) 160. (?) 58 Erbium Er 54.9 3 164.7 (?) 166. (?) 59 Thulium Tu 56 5 3 169.4 (?) 170.7(1) 60 Ytterbium Yt 57 2 3 171.7 (?) 173. (?) 61 Tantalum Ta 36 3 5 181.6 183. 69 Tungsten . . W 91 5 2 183.0 184.4 63 Osmium . . Os 47 4 4 189.6 191.0 64 Iridium Ir 47.9 4 191.6 193.0 65 Platinum ... . Pt 48.33 4 193.3 194.8 66 Gold Au 195.7 1 195.7 197.2 67 Mercury Hg 99 25 2 198.5 200.0 68 Thallium Tl 67.51 3 202.6 204.1 69 Lead . . Pb 102.67 2 205.34 206.9 70 Bismuth Bi 68.83 3 206.5 208.1 71 Thorium Th 57.7 4 230.8 232.6 72 IT 59.45 4 237.8 239.6 73 Helium He 74 Argon ... . /. . . A 19.8 (?) 75 Neon 9.5 (?) 17 INDEX The references are to marginal numbers. Part II is designated by the Roman numeral, and the appendix by the abbreviation Ap. The numbers are the same for the same topics in Parts I and II. Acetylene, 280, 281/ 2 . Acid defined, 30/V boric, 225. carbonic, 264. chloric, 541. chlorous, 541. hydrazoic, 343. hydrochloric, 533-535. hydrofluoric, 405, 406. hydrosulphuric, 510-512. hypochlorous, 538. hyponitrous, 314. nitric, 326-332. nitrohydrochloric, 331/ 2 , II. nitrous, 320. perchloric, 541. phosphoric, 494, 495. phosphorous, 493. salt defined, 30/ 6 . silicic, 473, 474. sulphuric, 519-526. sulphurous, 516. thiosulphuric, 522. Additive combination, 441. Additive compound denned, 336. Air, 338-395. Alkalies defined, 215. Allotropic change defined, 13. Alloy defined, 456. Alum, 461/j. Aluminium, 448-465. manufacture of, 462-465. Ammonia, 334-338. compounds, 336, 337. Amorphous defined, 21/ 6 . Analytic reaction defined, 81. Anhydrite, 578. Anhydrous defined, 21/ 7 . Anthracitic diamond, 240. Apatite, 485. Argon, 396. Arrhenius, 176. Asbestos, 443. Asymmetric defined, 185. Atmosphere, 388-395. ** ; Atom defined, 171. Atomic theory, 40/ 6 , 157-186. Atomic weights defined, 173. Atoms, heat capacity of, 175. Avogadro's hypothesis, 165. Azote. See Nitrogen, 306. Bacteria in nitrification, 328. in water, 377. Balances, Ap. 1. Barometer, Ap. 15 B. Base defined, 30/ 2 . Basic salt defined, 30/ 6 . Basicity of acids, table, Ap. 23. Bauxite, 449. 245 246 ELBMBNTAKY PRINCIPLES OF CHEMISTRY Bending glass tube, Ap. 9. Benzene, 282. Bergmann, 259. Berthollet, 37/ 6 . Beryllium. See Glucinum, 216-219. Berzelius, 466. Black, 259, 577. Blasting gelatin, 575. Bleaching powder, 547, 548. Blowpipe, use of, Ap. 21. Boiling denned, 11. Boiling point affected, 24/ l5 24/ 2 . apparatus, Ap. 17. denned, 24. Boiling temperature, constant of elevation, 130. constants, tables, 132-135. determination of, 127/j, 127/ 3 , II. molecular elevation of, 176. specific elevation of, 126. Bone black, 252. Borax, 228. bead, 228, II. Boric acid, 225. Boron, 220-228. Boyle, 66/3, 388. Boyle's Law, 66. apparatus for, 66, II. Brand, 484. Brimstone. See Sulphur. Bromine, 541/j. Bumping, 24/ 2 . Bunsen, 450. Butane, 278. Butylene, 279. Calcium, 577-584. Calorie denned, 50A, II. v' ?^-_ Carbon, 234-258. amorphous, 248-258. Carbon, gas, 251. Carbon as constituent, 247. dioxide, 259-268. dioxide, weight of one liter de- termined, 81/x, 81/2, II. monoxide, 272, 273. Carbonado, 240. Carbonates, 265, 269-271. Carbonic acid, 264. Carborundum, 258. Cavendish, 200, 326, 368, 388. Charcoal, 252, 254. Charles's Law, 67-67/ 4 . apparatus, 67, II. Chemical properties defined, 3. change, essential feature of, 13. change, secondary features, 14. Chemism, 51. Chemistry defined, 2. Chromite, 594. Chromium, 592-606. Chlorine, 527-532. manufacture, 54%-54"6. oxides, 536-540. Chloric acid. 541. Chlorous acid, 541. Chlorophyll, 266. Classification, early, 621, 622. Coal, 298-300. Coal gas, 301, 302. Cobalt, 592-606. Coefficient of expansion, 67, II. Coke, 253. Combining weights of elements, defined, 44, 45. determination of, 150-152. list of, 643, 644. weights of compounds, deter- mination of, 155. weights, system of, 61. Combustion, 34/ 4 , II, 354. INDEX 247 Combustion of sulphur, 13, II. Composition defined, 15. Compound defined, 27. Congelation defined, 10. Cooke, 623. Cordite, 576. Corundum, 449. Crookes, Sir William, 348. Cryolite, 401. Crystallization, 21/ 4 . water of, 21/ 7 . Cutting glass tube, Ap. 7. Dalton, 40, 159, 174. Dalton's Law (Charles's Law), 67/ 2 . Data, table, Ap. 22. Davy, Sir Humphry, 220, 235, 407, 442, 527, 549, 577. Decanting defined, 18/B, II. Decomposition defined and illus- trated, 16. D%epitation, !Q/ lt II. Dehy^ed defined, 21/ 7 . Deli^uesc^e^^efined, 21/ 8 . Destructive crafcyilation of wood, 303-305. Deville, 450. Dewar, 206, 395, 404. Diamond, 235-241. artificial, 241. Diffusibility, law of, 203/V Dimorphous defined, 21/ 6 . Dissociation, electrolytic, theory of, 176. Distillate defined, 12. Distillation defined, 12. Distilling apparatus, Ap. 18. Dolomite, 443. Dorcet, 235. Dulong, 339. Dulong and Petit, 100. law of, 95-106. Dumas, 621, 623. Dynamic chemistry defined, 26. Dynamite, 572. Ebullition defined, 11. Effervescence defined, 13/j (6), II. Efflorescent defined, 21/ 8 . Electric furnace, 241. Electrolysis defined, 202. Electrolytic dissociation theory, 176. Element defined, 27. Elements, free or combined in na- ture, 192. in the living organism, 191. list of, 643, 644. metallic, 193. non-metallic, 193. their chemical activity, 198. their distribution, 189. their oxidation heat, its range, 199. their range in boiling point, 195. their range in density, 196. their relative quantity, 190. Energy, chemical, familiar forms, 60. chemical and electric, 59. defined, 1. persistence or conservation of, 34/ 9 . persistence or conservation of, in chemical phenomena, 50-60. Endothermic defined, 54. Epsom salt, 443. Equations, 63-63/ 2 . Equivalent proportions, law of, 41-46. weight defined, 42-44. weights, accuracy of, 145-148. weights, basis of, 146. 248 ELEMENTARY PRINCIPLES OF CHEMISTRY Equivalent weights, determination of, 142-149. weights, list of, 643, 644. Ethane, 278. Ethylene, 279. Evaporating to dryness, Ap. 12. Evaporation defined, 11. Exchange, double, 18. Exothermic defined, 54. Explosives, 555-576. Factors defined, 26. Feldspar, 482. Fertilizers (phosphates), 502. Filtration, filtrate defined, 18 A (c), II. Filtration, the manipulation, Ap. 13. Fixed proportions, the law of, 37- 37/ 8 . Flame, nature of, 284-292, II. Fluor spar, 401. Fluorine, 401-406. Fluorite, 401. Formula of compounds, determi- nation, 153-155. Fornmla weight, 62/^. Freezing defined, 10. Freezing point defined, 23. constants of depression, 117. constants of depression, table, 120-123. depression of, 23/i. molecular depression of, 176. specific depression of, 110/ B . specific depression of, determi- nation, 111/, .-111/4, II. Fusion defined, 10. Gallium, 638. Galvanic cell, 59. Gas, manipulation of, A p. 19. Gas, volumetric proportions, 47- 47/7. Gases, kinetic theory of, 164. Gases, volume, pressure, and tem- perature, 66-67/ 4 . Gay-Lussac, 47/ 7 , 89, 220. Gay-Lussac's Law (specific gravi- ties), 71-94. see Charles's Law, 67/ 2 . (volumetric proportions), 47. Generator, gas, A p. 6. Germanium, 638. Gladstone, 625. Glass, 477-481. soluble, 476. Glucinum, 216-219. Graduated cylinder, Ap. 14. Graphite, 243-246. Guano, 502. Guncotton, 573-576. Gunpowder, 556-564. Gypsum, 507, 578. _> Halogens defined. 541/i. Hardness of water, 375-375/ 3 . Heat capacity of atoms, 175. Heat disturbance, 50. Heat of formation defined, 53. Heat, measurement of, 54, 55, 55/j . of neutralization, 50/ 4 , II. summation, law of, 55. unit of, defined, 50A, II. Heats of formation, list, 57. Heating a crucible, Ap. 11. a test-tube, Ap. 2. Helium, 400. Helmholtz, 50, note. Hematite, 594. Hess, 55. Homologous defined, 278. Humboldt, 47/ 7 . Hydrazine, 341. INDEX 249 Hydrazoic acid, 343. Hydrocarbons, 274-283. Hydrochloric acid, 533-535. Hydrofluoric acid, 405-406. Hydrogen, 200-209. dioxide, 366, 367. peroxide, 366, 367. Hydrosulphuric acid, 510-512. Hydroxylamine, 342. Hygroscopic defined, 21 / 8 . Hypochlorous acid, 538. Hyponitrous acid, 314. Hyposulphite. See Thiosulphate, 522. Increment of volume, 67, II. Interaction defined, 26. Iodine, 541A. lonization, 176. Iron, 592-606. commercial, 607-619. Isomerism, stereo or physical, 183- 185. Isomers defined, 179. Isomorphous defined, 21/ 6 . Joule, 50, note. Kaolin, 482. Kelvin, Lord, 162. Kinetic theory of gases, 164. Kopp and Neumann, law of, 103, 104. Krypton, 399. Lake (as to dyeing) defined, 460. Lampblack, 250. Lavoisier, 34/ B , 37/ 6 , 200, 235, 259, 306, 326, 388. Le Bel, 185. Lecoq de Boisbaudran, 638. Lime, 581-583. Lime, hydraulic, 587. preparation, 585. Limonite, 594. Lithium, 210-214. Magnesite, 443. Magnesium, 442-447. Magnetite, 594. Manganese, 592-606. Mariotte, 66/ 3 . Marsh-gas. See Methane, 275-277. Mass defined, 34. persistence or conservation of, 34/6- Matches, 505. Matter defined, 1. Mayer, J. A., 50, note. Melting defined, 10. Melting-point apparatus, Ap. 16. defined, 23. Mendeleeff, 625, 638, 639, 641. Metal defined, 29. Metallurgy defined, 257. Metamers defined, 180. Metargon, 399. Metathetic reaction defined, 31. Methane, 275-277. Meyer, Lothar, 625. Mitscherlich, 108. Law of, 107, 108. Mixture defined, 28. . Moissan, 223, 241, 256, 401, 579, 580. Molecular depression of freezing point, 176. Molecular elevation of boiling tem- perature, 176. Molecular theory, 160-163. weights defined, 167. Molecule defined, 166. of compounds, divisible, 168. of elements, divisible, 160, 170: 250 ELEMENTARY PRINCIPLES OF CHEMISTRY Molecules, estimated size, 162. structure of, 177-182. Monobasic defined, 333. Mordant (as to dyeing) defined, 460. Morley, 147. Mortar, 586. Mortar and pestle, Ap. 5. Multiple proportions, law of, 40- 40/ 8 . Nascent state defined, 208. theoretic explanation of, 208/ 2 . Neon, 398. Neumann. See Kopp. Neutral, neutralized defined, 30/ 4 . Neutralization, heat of, 50/ 4 , II. Neutralizing, the operation of, 37, II. Newlands, 625. Nickel, 592-606. Nilson, 638. Nitrates, 333. Nitric acid, 326-332. Nitric oxide. See Dioxide, 317, 318. Nitrides, 340. Nitrification, 328. Nitrogen, 306-311. chloride, 339. dioxide, 317, 318. monoxide, 312-316. oxides 1 ^ 311. pentoxide, 325. peroxide, 323. tetroxide, 322-324. trioxide, 319-321. Nitroglycerine, 569-572. Nitrohydrochloric acid, 331/ 2 , II. Nitrous acid, 320. Nitrous oxide. See Monoxide, 312- 316. Nomenclature, 65-65/ c . Non-metal, 29. Normal salt defined, 30/ 6 . Notation for compounds, 62-62/ 2 . for\elements, 62. preliminary statement, 19, II. Occlusion defined, 205. Octaves, law of, 625. Olefins, 279. Organic chemistry defined, 283. Oxidation (as to valence), 441/. Oxides, classes of, 355. Oxidizing agents defined, 48, II. flame, Ap. 21. Oxygen, 350-357. weight of one liter, 71 A, II. Ozone, 358-365. Paracelsus, 200. Paraffins, 278. Pasteur, 184. Perchloric acid, 541. Periodicity, 589, 620, 625, 640. Permanganate, 604. Pestle and mortar, Ap. 5. Petit. See Dulong. Petroleum, 293-297. Phosphine, 496. Phosphoric acid, 494, 495. Phosphorous acid, 493. Phosphorus, 484-501. manufacture, 498-501. Physical isomerism, 183-185. properties defined, 3. properties enumerated, 4-12. Physico-chemical properties de- fined, 3. Physics defined, 2. Plaster of Paris, 588. -'olymerization defined, 281/ 2 . 5 olymers defined, 179. polymorphous defined, 21/ 6 . 'orcelain, 482. INDEX 251 Potassium, 549-554. Pouring from reagent bottle, Ap. 3. Precipitate, precipitation defined, 18/C, II. Priestley, 306, 312, 334, 350, 388. Products defined, 26. Propane, 278. Propylene, 279. Protoplasm, its composition, 191. Proust, 37/ 6 , 40. Prout's hypothesis, 624. Proximate analysis defined, 26. Pyrites, 507, 594. Pyrolusite, 594. Qualitative analysis defined, 26. Quantitative analysis defined, 26. Ramsay, 396, 398-400. Raoult, law of (boiling tempera- ture), 124-135. (freezing point), 109-123. Rayleigh, Lord, 396. Reactions classified, 31. defined, 26. Reducing agents defined, 48, II. flame, Ap. 21. Reduction (as to valence), 441/a. Richter, 40, 41/ 9 . Ruby, 449. Rutherford, 306. Salt defined, 30/ 3 . Sapphire, 449. Saturated defined (as to valence), 441. Saturation defined (as to solution), 21/3- Scandium, 638. Scheele, 350, 484, 527. Siderite, 594. Silicates, some uses of, 476-483. Silicic acid, 473, 474. Silicon, 466-475. Soapstone. 443. Sodium, 407-410. carbonate, 413. chloride, 411, 412. hydroxide, 419, 420. Solubility of salts, list, Ap. 24. Solute defined, 21. Solution defined, 21. heat of, 22. Solvent defined, 21. Specific depression of freezing point defined, 110/a. determination, lll/i-111/4, II. Specific elevation of boiling tem- perature defined, 126. determination, 127/ 1 -127/ 5 , II. Specific gravity defined, 7. determination, 8. Specific heat defined, of compounds, 1 determination, 95/ 2 , 95/ 3 , II. of elements, table, 105. law of, 95-106. Stas, 145. Static chemistry defined, 26. Steel, 607. Stereo-isomerism, 183-185. Stoichiometry, 64, 64/j. Stoneware, 483. Structure of molecules, 177-182. Sublimate, sublimation, defined, 12. Substances defined, 1. Substitution defined, 17, 31. Sulphur, 507-509/4. oxides, 514-518. Sulphuric acid, 519-526. Sulphurous acid, 516. Superphosphate, 503. Surfusion, 23. 252 ELEMENTARY PRINCIPLES OF CHEMISTRY Survey, general, first nine ele- ments, 422-433. general, seventeen elements, 589, 590. Synthetic reaction defined, 31. System, chemical, defined, 26. Talc, 443. Thenard, 220. Theory, atomic, 157-186. of electrolytic dissociation, 176. kinetic, of gases, 164. Thermometer, Ap. 15 A. Thiosulphuric acid, 522. Toluene, 282. Ultimate analysis defined, 26. Valence, 434-441 /.. Van't Hoff, 185. Vapor-densities, table, 93, 94. Vapor-density, data, discussion, 72-87. defined, 71. determination, 71/A, 71/x, II. effect of temperature on, 88, 89. law of, 71-94. Volumetric proportions, 47-47/ 7 . Water, 368-370. action on lead and zinc, 387. for drinking, 377-377/ 5 . purification of, 378-386. Waters, natural, 371-377/ 6 . Weighing, Ap. 1. Weights, Ap. 1. Wenzel, 41/ 9 . Winkler, 638. Wohler, 178, 216, 450. Wood, destructive distillation of, 303-305. Xenon, 399. PART II EXPERIMENTAL ILLUSTRATIONS KECOMMENDATIONS AS TO NOTES THE making of notes is a very important item in the study of a subject in the laboratory. The student should early in his course form the habit of writing them care- fully, thoughtfully, and systematically. The aim should be to describe the essential features of method and of appa- ratus, then the things observed, then the things learned from the experiment; for every experiment is designed to teach something. The aim should also be to make the description clear, exact, and simple. The following details of plan are recommended : To use a book about six by nine inches in size, one whose leaves will lie flat when opened, with plain, unruled pages ; to enter notes only on the right- hand page, leaving the left-hand for topics, corrections, ad- ditions, and the numerical work of calculations ; to enter always the particular topic which the experiment is to illus- trate ; as a minor item, to enter the date of work ; for con- venience, to have the name of the owner on the outside of the cover. The notes should be written in final form in the laboratory, and not copied. It is particularly urged that the student should read through the directions for an experiment before starting upon its performance, and that the experiment should pre- cede the consideration of the corresponding topic in Part I. THE ELEMENTARY PRINCIPLES OF CHEMISTRY PART II EXPERIMENTAL ILLUSTRATIONS CHAPTER I INTRODUCTION Read Part I, Nos. 1-6, before beginning the experiments. 1. Physical Properties of Sulphur Observe as to odor, color, hardness, form (crystalline 1-6* that is, regular geometric form). To determine its specific gravity: Weigh carefully to 7 tenths about 10 grams of dry sulphur, in small lumps, free of dust. Weigh a test-tube, filled with water and corked. Put the sulphur in the test-tube, refill with water, cork, and weigh (see Appendix, 1). The specific gravity of sulphur equals the weight of sulphur divided by the weight of an equal volume of water. Calculate the specific gravity from the observations made. Suggest another method for solids ; also one suitable for solids 8 which are soluble in water ; a method for liquids ; for gases. Electrification. Rub a large lump briskly on the dry 9 towel, or, better, on some woolen stuff. It becomes electri- fied and capable of attracting particles, like bits of paper. * The marginal numbers are the same for the same topics in both Part I and Part II. 1 V 2 ELEMENTARY PRINCIPLES OF CHEMISTRY 10 Effects of heating. Heat in a dry test-tube some lumps of dry sulphur, the test-tube being about half filled. It is well to make a holder by folding a strip of paper once or twice upon itself, to wrap this about the test-tube and to seize the ends close to the tube with the fingers, or ; better, with the tongs (see Appendix, 2). Apply the heat grad- ually by passing the tube slowly back and forth through the flame. Have a vessel filled with water on the table close at hand. Should the tube crack and the sulphur take fire, hold quickly over the water. Let the temperature rise slowly while the sulphur is melting. 2s ote the color. Keep the surface of the liquid in gentle motion ; note the partial 1 1-12 solidification. Continuing the heat, note the second lique- faction, the beginning of boiling, and the white dustlike deposit (sublimate). Now, as the sulphur is beginning to boil, move the burner to one side ; carefully and slowly turn the melted sulphur into the water, moving the tube to and fro so as to string it out. Set aside the tube with the small residue of sulphur and observe it from time to time as it cools (solidification and crystallization}. Examine the por- tion which has been suddenly cooled in the water (allotro- pism), comparing it with the original. Select a good sample of it, and, without lumping it together, dry it by pressing and rubbing first between folds of the towel, then between absorbent papers (filter papers). Weigh the sample care- fully and leave it for twenty-four hours. Does it change in weight in passing from the plastic amorphous (see Part I, No. 21/6 ) to the brittle crystalline form ? 12/1 Crystallization from fusion (see No. 21/ 4 , Part I). Fuse some dry sulphur in a crucible, keeping the temperature as low as will suffice to melt ; let it cool slowly until a crust has formed ; puncture the crust and pour out the remaining liquid; break the crucible and note the appearance of the crystals known as monoclinic sulphur ; examine them twenty-four hours later. (It is recommended that the instructor do this experiment for the class.) 12/2 Crystallization from solution (see No. 21/ 4 , Part I). Dissolve some sulphur in carbon disulphide to saturation, or nearly so, decant (that is, INTRODUCTION 3 drain off) or filter the liquid and allow it to evaporate in the hood. Examine the crystals known as orthorhombic sulphur ; compare with those of the preceding experiment (dimorphism}. Do they change on standing ? As carbon disulphide is very volatile and combustible, great care should be taken to have no flame near. (It is recommended that the instructor do this. experiment for the class.) 2. Chemical Properties of Sulphur Burning or combustion. Ignite a small lump of sulphur 13 (half the size of a pea) on a suitable surface (an earthen saucer serves well). Describe the phenomenon. The gase- ous product is sulphur dioxide ; observe its odor cautiously, and its effect on litmus conveniently done by sticking a piece of wet, blue litmus paper to the bottom of a beaker and inverting the latter over the burning sulphur. Com- pare the sulphur dioxide in these respects with the sulphur. If you have a sample of the plastic sulphur, burn and test this in the same manner. EXPLANATORY NOTE. An invisible gas, namely, oxygen, a constituent of the atmosphere, takes part in this change, combining with the sulphur and forming the sulphur dioxide, which is quite different from both the sulphur and the oxygen. Behavior with zinc. Weigh out carefully 3.20 grams of 13/1 finely divided sulphur (sulphur sublimate, commonly called sulphur flowers) ; also 6.50 grams of finely divided zinc (zinc dust). Mix these thoroughly in a mortar, so that the original powders can not be distinguished (see Appen- dix, 5). Take out a small portion, about as much as could be heaped on a five-cent piece (putting the remainder well to one side) ; heat this strongly on the earthen saucer or on the asbestos board by turning the gas flame down upon it, or, still better, on the spatula blade by thrusting it into the flame. Do this cautiously, as it may flash up somewhat like gunpowder. In experiments like this, the face should never be held over the material Describe the phenomenon, note the ashlike product, compare it with the original sub- 4 ELEMENTARY PRINCIPLES OF CHEMISTRY stances, sulphur and zinc, by inspection. Does it melt and burn like sulphur ? (a) Put a little of the mixed powders from the mortar in a test-tube, half fill with water, then, closing the tube with the thumb, shake it vigorously, and finally let the powder settle (see Appendix, 4). Does the powder dissolve ? Do you see the sulphur tend to separate from the mixture with the zinc ? In a similar manner shake some of the ash- like product with water. Does it dissolve ? Do you see any tendency of the sulphur to separate from the zinc ? Does zinc dissolve in water ? Does sulphur dissolve in water ? (b) Put a small quantity of the ashlike product in a test-tube and add to it a few drops of hydrochloric acid, better designated as hydrogen chloride (see Appendix, 3). Note the phenomenon of effervescence (bubbles of gas in the liquid ; how does it differ from boiling ?), indicating the formation or liberation of a gas not previously apparent. Note the odor of this gas (with care, for the substance is very poisonous) and its effect on a bit of filter paper wet with a solution of lead acetate, and held at the mouth of the test- tube. Does the ashlike product disappear in the acid ? If not, warm the tube slightly. Compare this with the behavior of a little zinc treated in another test-tube with hydrogen chloride in a similar way (does the zinc powder disappear ?) ; also of sulphur thus treated ; also of the mixed zinc and sulphur (does this powder disappear ?). EXPLANATORY NOTE. In this experiment you have brought about a chemical change. Starting with zinc and sulphur, mixing them inti- mately and applying heat, you have caused the zinc and the sulphur to disappear, and in their place you have a substance which is distinctly different from each of the original substances. This is a change of identity, a chemical change ; and with this meaning, the new substance, which for the moment may be designated x, is said to contain zinc and sulphur, and the latter are said to be combined in the substance x. To the question, Does the substance x contain nothing but zinc and sul- phur ? the answer will be suggested by the next experiment. INTRODUCTION 5 Behavior with iron. Weigh out carefully 3.20 grams of 13/2 sulphur powder ; also 5.60 grams of finely divided iron (iron dust). Mix them thoroughly in the mortar. Take out a portion, ignite it, and compare the product with the iron, with the sulphur, and with the mixture by inspection and by its behavior with hydrogen chloride, as in the preceding experiment. Does the gas which is obtained have the same odor and action on paper wet with lead acetate as in the preceding ? (a) Vary this procedure as follows : Put a portion of the mixed iron and sulphur in a dry test-tube, and cork the lat- ter ; weigh it and its contents carefully to tenths ; then, holding it nearly horizontal, tap it gently so that the pow- der spreads out in an elongated pile from the bottom of the tube ; let the cork lie loosely in the mouth of the test-tube, and heat the other end in the flame just enough to start the action ; then remove from the flame. When the action has ceased, let cool, and weigh again with contents. With a lit- tle care, you will be able to do this without losing anything from the tube during the operation. Save the product. EXPLANATORY NOTE. Again you have a chemical change. The substance produced (y) contains iron and sulphur as constituents, but is neither iron nor sulphur ; moreover, it contains nothing bid these, since what you have in the tube after the change weighs no more than the iron and the sulphur which you put in. Iron and sulphur then must be the sole constituents of the substance y ; zinc and sulphur are likewise the sole constituents of the substance x. The first is therefore named iron sulphide, and is said to be composed of iron and sulphur, and the second is zinc sulphide, and is composed of zinc and sulphur. The " changes by which these are produced fall under a common type, which may receive a general expression, thus : Substance A and substance B became substance A B (composition) ; or, to express it more concisely, A + B = A B. But you have further illustration of chemical change in the produc- tion of a peculiar gas by the action of the hydrogen chloride on both zinc sulphide and iron sulphide. This gas is clearly very different from all these other substances. You have noted the same odor and the ' 18 6 ELEMENTARY PRINCIPLES OF CHEMISTRY same action on the lead paper in both cases, and, in fact, the same gas is produced by both substances. If this is so, evidently neither iron nor zinc can enter into its composition. But the sulphur of one or of the other does so enter, yet the gas is markedly different from sulphur itself and from the gas produced by burning sulphur (see Exp. 13). The sole other constituent of this peculiar gas is hydrogen, which comes from the hydrogen chloride, of which also it is a constituent. The gas is therefore named hydrogen sulphide. It is also known as hydrosul- phuric acid. 13/3 To show that hydrogen sulphide contains sulphur: Place about 1 gram (not more than 2) of the iron sulphide, made in the previous experiment, in the gas generator (see Ap- pendix, 6-9) with just enough water to seal the thistle-tube. Drop the delivery tube into a test-tube about one third filled with nitric acid (handle this substance with care, for it is very corrosive *), so that the gas shall bubble through this liquid. Pour a few drops of hydrogen chloride into the thistle-tube of the generator, so as to produce a quick stream of bubbles through the liquid. Place a bit of paper wet with lead acetate over the mouth of the test-tube. As the gas from the generator passes through the nitric acid you see small particles of white or yellowish substance appearing, which, when the action has ceased, or nearly so, will have collected into a mass from which you may rinse the acid with water, and which you may easily recognize as sulphur by inspection and by burning. At the same time the lead paper may, perhaps, show no blackening, since the sulphur is produced only at the expense of the hydrogen - sulphide, although enough of the latter to stain the paper may escape destruction by the nitric acid. EXPLANATORY NOTE. Hydrogen sulphide, therefore, contains sul- phur, and, whenever this gas is produced, it may be taken as evidence of sulphur as a constituent in the substances from which it is produced. * If nitric acid is spilled on the skin, clothing, or table, the spot should be as quickly as possible moistened with ammonia (side-table), : and then rinsed with water. INTRODUCTION 7 QUESTIONS. How can you differentiate (that is, distinguish by dif- ferences) between sulphur, and oxygen, and sulphur dioxide ? Between zinc, and sulphur, and the mixture of zinc and sulphur, and zinc sul- phide? Between iron, and sulphur, and the mixture of iron and sul- phur, and iron sulphide? What constitutes the difference, by defi- nition, between a mixture and a chemical compound ? (see Nos. 13 and 28, Part I.) How do you bring about the chemical change between sulphur and air, between zinc and sulphur, between iron and sulphur? What evidence by observation have you that heat is produced by these three changes? Does zinc sulphide, made by heating zinc and sulphur, change back to zinc and sulphur by cooling? Does iron sulphide re- verse the change by cooling ? Is the change from iron and sulphur mixed to iron sulphide accompanied by change in the total weight? (see No. 14, Part I.) 3. Additional Illustrations of Chemical Change Composition. Put in a dry test-tube about 1 gram of lead 15/1 dust and 0.2 of a gram of iodine (iodine must be weighed on glass or paper, as it corrodes metal) ; cork the tube and weigh it with its contents carefully ; let the cork lie loosely in the mouth of the test-tube and warm the mixture very slightly, at first barely touching the flame with the bottom of the tube. When the first action is over, the sub- stance may be heated until it melts. A little of the iodine may escape the change, making a purple vapor and a dark deposit on the upper part of the tube. When the tube is cool, weigh it with its content. Describe the phenomenon, and compare the yellow substance with the lead and with the iodine. Preserve a sample of it for future comparison. EXPLANATORY NOTE. From lead and iodine is produced a substance which does not give off a purple vapor, nor impart a brown color to water, as does iodine, and which is clearly neither iodine nor lead ; yet by the test of weight it can contain nothing more than iodine and lead. This bright-yellow substance is named lead iodide. Does it dissolve in water? Composition. Put about 0.3 of a gram of magnesium 15/2 ribbon, or powder, in the porcelain crucible (see Appen- dix, 11). Weigh carefully the crucible with its content 8 ELEMENTARY PRINCIPLES OF CHEMISTRY and its lid. Support the same on the pipe-stem triangle, and the latter on the iron ring of the stand. Apply the full heat of the flame to the bottom of the crucible for ten or fifteen minutes. Eaise the lid slightly to observe what is taking place, but do not let any of the white smoke escape. During the last few minutes the lid should be entirely removed. Let the crucible cool, and weigh it again with content and lid. How do the two weights com- pare ? Describe the phenomenon and compare the sub- stance after heating with the original.* EXPLANATORY NOTE. The bright metallic ribbon or powder is changed, simply by heating in contact with the air, to a very different substance a grayish chalklike powder ; but this weighs more than the original. It must be, therefore, that something besides magnesium is contained in the product. This " something " comes from the air, is an invisible gas, and is named oxygen. The white powder is named mag- nesium oxide. Try also heating a short piece of the ribbon directly in the flame, holding it with the tongs ; likewise a little of the oxide. 15/3 Composition. In a similar manner heat in the crucible about 2 grams of zinc dust, weighing the crucible and contents before and after the ignition. Compare the weights and describe the phenomenon. Zinc oxide is thus formed yellow while hot, white when cold. Try also heat- ing a little of the zinc dust on the spatula blade. Preserve a little of the zinc oxide for future comparison. 15/4 Composition. Spread about 4 grams of lead dust on the crucible lid, weighing the whole ; then ignite over the flame for fifteen or twenty minutes and weigh again. The red- dish-yellow powder is lead oxide. Preserve a little of this. EXPLANATORY NOTE. The statement made concerning magnesium oxide applies also to zinc and lead oxides. These four instances of composition are changes of the general form: Substance A and substance B become substance A B ; or A + B = A B. * Clean the crucible by using a few drops of hydrochloric acid, warm- ing gently, then rinsing with water. Scouring with sand may be helpful. INTRODUCTION 9 QUESTION. Have you evidence that, when magnesium changes to magnesium oxide, and zinc to zinc oxide, heat is produced f Decomposition. Take a quantity of lead nitrate, about 16/1 1 gram (twice the size of a pea), pulverize finely in a clean, dry mortar, shake the powder into a clean, dry test- tube, wipe off any particles adhering to the upper part of the tube, cork, and weigh the tube and contents. Then, holding the tube nearly horizontal (paper holder), with the cork loosely placed in its mouth, warm the substance gently until the snapping (decrepitation) ceases and the powder melts quietly ; remove the cork and continue the heating until the bubbling ceases and the glass begins to soften. Note the effect of the brown gas on wet, blue litmus paper. Let cool, and weigh the tube and contents, including the cork. Break the tube and examine the reddish-yellow resi- due, lead oxide. Save a little of this for future reference. QUESTIONS. Does the residue dissolve in water? Does the lead nitrate dissolve in water ? Does the brown gas condense at all on the upper part of the test-tube as you saw in the heating of sulphur f (see Exp. 11.) How does the weight of the residue compare with that of the lead nitrate ? EXPLANATORY NOTE. In this experiment you see one substance yielding at least two others, distinctly differing from each other as well as from the original ; and at least one of these that is, the residue in the tube weighs less than the original, which has disappeared, although none of the latter, unchanged, has passed out of the tube. The gas which has passed out, named nitrogen tetroxide, and the residue, named lead oxide (see Exp. 15/ 4 ), which is left in the tube, must have been contained in the lead nitrate which has been destroyed ; the latter has been broken up into at least two other substances, each of which weighs less than the original. Such a change is called decomposition. It may be described under the general type : Substance A B becomes sub- stance A and substance B ; or A B = A + B. Decomposition. Place about 1 gram of zinc nitrate in a 16/2 test-tube, weigh the tube and contents, then heat. Describe the phenomenon, and note the clear liquid which condenses 10 ELEMENTARY PRINCIPLES OF CHEMISTRY on the upper part of the tube. This is water ; it is slowly driven from the tube by heat. Later, note the brown gas, which appears as in the preceding experiment. This is nitrogen tetroxide, and it in turn is driven from the tube. Finally, there is left a powder which no longer melts, and which is yellow while hot and white when cold. This is zinc oxide. Weigh again the tube and contents. Examine the residue left in the tube (comparing with the substance in Exp. 15/ 3 ), and save a portion for future tests. QUESTIONS. Does the residue weigh more or less than the zinc nitrate from which it comes f Does it dissolve in water ? Does zinc nitrate dissolve in water ? EXPLANATORY NOTE. You see the substance, zinc nitrate, destroyed by heat, yielding at least three distinctly different substances water, nitrogen tetroxide, and zinc oxide ; and the last of these, and by infer- ence each of the others also, weighs less than the original substance, and all of them must have been contained in the original zinc nitrate. This change falls under the same type as the preceding : 17/1 Substitution. (a) Place a small quantity of granulated zinc, about 0.5 of a gram, in a test-tube ; add to it a few drops of hydrochloric acid (hydrogen chloride). Note the effervescence which implies the liberation of a gas. Close the tube with the thumb for a few seconds, then open it, holding its mouth to the gas flame. Note what takes place. EXPLANATORY NOTE. The colorless invisible gas liberated in this change burns with an almost invisible flame, and makes with air an explosive mixture. This causes the slight sound when the content of the tube is ignited. Larger quantities may be dangerously explosive, and this fact should always be in mind when dealing with this gas which is named hydrogen. (b) Next, place 1 gram of granulated zinc in a small evaporating dish which, with a small glass stirring-rod (see Appendix, 12), has been previously weighed (to tenths is sufficient). Add about one quarter of a test-tubeful of hydrochloric acid, in portions at a time, warming some- INTRODUCTION 11 what the contents of the dish. Holding a lighted match at the surface of the liquid while the bubbles of hydrogen are breaking, will again show the combustibility of this gas. When the zinc has disappeared except a few black specks, and the bubbles of gas no longer appear, increase the heat somewhat, holding the burner in the hand and touching the tip of the small flame to the bottom of the dish from time to time, as may be needed to keep the liquid quietly boiling (see Appendix, 12). This soon thickens and de- posits at the edges. Continue the heating, avoid spatter- ing as much as possible, and soon white fumes appear, and the substance solidifies if cooled. At this point cease the heating and, when the dish is cool, weigh it with its con- tents. Describe the substance in the dish. Does it weigh more or less than the zinc taken ? Let some of it remain exposed to the atmosphere for a short time and note the change. EXPLANATORY NOTE. Hydrochloric acid contains as sole constit- uents hydrogen and chlorine, the material which you use being this substance dissolved in water. When this is brought in contact with zinc a change takes place, and, when this is complete and the liquid boiled to dryness, the product just seen is the result. Now this weighs more than the zinc, nearly double, and, therefore, must contain some- thing besides zinc ; yet it can not contain anything more than zinc, hydrochloric acid, and water ; indeed, it can not contain all of these, for hydrogen has passed into the air, as you have seen, and the water also has been boiled away. In fact, the residue in the dish is a sub- stance which contains as sole constituents the zinc used and the chlorine previously contained in the acid, although it is very different from each of these. It is named zinc chloride. Exposure of this to the air for a short time will show you incidentally one of its proper- ties; it absorbs water from the atmosphere and liquefies in conse- quence (deliquescence). In this change it is seen that a constituent of hydrochloric acid namely, hydrogen leaves the substance, and zinc may be said to take its place ; so that, starting with zinc and hydrogen chloride, you obtain hydrogen and zinc chloride. Such a change is called substitution. It may be represented by the general form : 12 ELEMENTARY PRINCIPLES OF CHEMISTRY 17/2 Substitution. Optional experiment. (a) Place about 0.5 of a gram of granulated zinc in a test-tube, add dilute sul- phuric acid, warm, and test the gas at the flame as in the preceding experiment. (b) Next place about 3 grams of granulated zinc in the evaporating dish, add about one t'est-tubeful of dilute sul- phuric acid (one volume of concentrated acid to four of water) and an equal volume of water, and warm. When the zinc has disappeared, or nearly so, filter (see Appendix, 13), boil down the liquid somewhat, and set aside to crys- tallize. Crystals of zinc sulphate are thus obtained. The relation of this substance to sulphuric acid is similar to that of zinc chloride to hydrochloric acid.* 17/3 Substitution. Place about 3 grams of lead powder in a small evaporating dish, add about one quarter test-tubeful of nitric acid and an equal volume of water. Heat gently until the action ceases, filter the liquid, and set it aside to crystallize. When there is a considerable deposit of crys- tals, drain off the liquid, and throw the crystals on two or three thicknesses of filter paper to absorb the adhering liquid. These are crystals of lead nitrate. Heat a por- tion of these and compare the result with that in Kos. 15/4, and 16/j. Save a portion of the crystals, labeled, for later use. .17/4 Substitution. Optional experiment. Dissolve about 10 grams of crystallized copper sulphate in about 50 cubic centimeters (see Appendix, 14) of boiling water, using the larger evaporating dish. Add to this 4 grams of granu- lated zinc, and boil the liquid gently until the blue color has entirely disappeared. Filter; concentrate the liquid to about one third its volume ; filter again if necessary, and set aside to crystallize. What is the reddish-brown powder * The irritation which may be caused by the fumes and spray in this experiment, if many are working it at the same time, may be re- lieved by inhaling moderately the fumes from the bottle of ammonium hydroxide (side-table). OF THE INTRODUCTION UN1VEH 13 which appears when the zinc is added ? Adt&ps of nitric acid to this powder after filtration. What is the substance which crystallizes in the filtrate ? EXPLANATORY NOTE. Copper sulphate bears the same relation to copper and sulphuric acid that zinc sulphate bears to zinc and sul- phuric acid (see Exp. 17/ 2 ) that is, copper has taken the place of hy- drogen as a constituent in sulphuric acid. When zinc is brought in contact with the solution of copper sulphate, the copper, in turn, is dis- placed, appearing as a powder, which, with nitric acid, reproduces a blue solution ; while the zinc, with the sulphuric acid, forms zinc sulphate, crystallizing as seen in Exp. 17/ 2 . Iron may likewise thus substitute itself for copper as constituent. If the action is allowed to take place slowly in dilute solution, the copper may be deposited as a film on an object of iron, such as a knife-blade. This is cajled " plating." Double exchange. The purpose of this experiment is to 18 illustrate a reaction of the type : This is known as double exchange or double decomposition ; it is also described as a metathetic reaction. The experi- ment is rather long, so its plan is here given as a whole : A. To prepare zinc iodide from zinc and iodine : (a) To show that this is neither zinc nor iodine. (b) To show that it contains both zinc and iodine. (c) To prepare the solution of zinc iodide in water. B. To prepare lime sulphide in solution by water from lime and sulphur. .(d) To show that this contains sulphur as con- stituent. C. To mix the solution of zinc iodide and that of lime sulphide, a chemical change taking place by which a liquid and an insoluble solid are produced. (e) To show that the liquid now contains lime and iodine in combination, but no zinc nor sul- phur. 14 ELEMENTARY PRINCIPLES OF CHEMISTRY (/) To show that the insoluble solid contains zinc and sulphur in combination, but no iodine nor lime. Conclusion. Zinc iodide and lime sulphide have become zinc sulphide and lime (calcium) iodide. 18/A A. To prepare zinc iodide. AVeigh a clean, small evapo- rating dish and a short glass rod to tenths of a gram. Place in the dish 1 gram of granulated zinc and a few drops of water. Weigh out 3.9 grams of iodine (on glass or on paper). Add about one fourth or one fifth of the iodine to the zinc in the dish, stir, and warm very slightly, just enough to start the change which itself produces much heat. When the action seems to cease, add again about the same quan- tity of iodine, stirring constantly, and so on until all has been used. This gives a small volume of dark-colored liquid. Holding the burner in the hand (see Appendix, 12), heat from time to time to keep the liquid gently boil- ing. There is some escape of purple fumes for a few sec- onds ; this vapor is irritating, so avoid inhaling it. When the contents thicken there is a tendency to spatter ; pre- vent this by constant stirring and moderating the heat. Finally, when the contents seem thoroughly dry, white fumes begin to escape. At this point cease heating. The product is a dry powder slightly brownish in color. When the dish is cool, weigh it with its contents. The substance weighs more than the zinc, but somewhat less than the zinc and iodine together, owing to loss by the heating.* (a) To show that this is neither zinc nor iodine: You have seen that it shows no purple fumes upon heating, as iodine would show. Place a small portion, such as would be taken on the point of a penknife, in a test-tube with a few drops of * Iodine stains may generally ba. removed by using sodium sulphite or ammonium sulphide. INTRODUCTION 15 water. It dissolves as zinc would not, but with no brown color such as iodine would impart. (b) To show that it contains both zinc and iodine : Place another small portion in a small evaporating dish, add to it a drop of nitric acid, and heat to complete dryness. You see the purple vapor of iodine, and have left in the dish ' zinc oxide, as seen in Exps. 15/ 3 and 16/ 2 . (c) Dissolve the rest of the product, zinc iodide, in 50 cubic centimeters of water (see Appendix, 14). You may see some of the original zinc unacted upon. Filter the solution (see Appendix, 13). Turn a few drops of the filtrate (that is, the liquid which has passed through the filter) into a test-tube, and add a few particles of lead nitrate (see Exp. 17/ 3 ) previously dissolved in water. The bright- yellow powder now appearing in the liquid is lead iodide, seen also in Exp. 15/j, and must be taken as further evi- dence of iodine as a constituent in the substance formed in 18/A. Set aside the remainder of the solution and label it zinc iodide. To what type of change does this formation of zinc iodide belong ? B. To prepare calcium or lime sulphide in solution. 18/B Place 5 grams, roughly weighed, of powdered lime in the larger evaporating dish. Add 100 cubic centimeters of water and boil. Touch a piece of red litmus paper to the liquid and note the effect. Add to the contents of the dish 3.2 grams, roughly weighed, of sulphur pow- der (flowers) and boil, stirring constantly. Keep the liquid boiling gently for a few minutes, until it becomes colored quite a deep yellow. Cease the heating, and the powder quickly settles to the bottom. Drain off the liquid as much as practicable without disturbing the sediment (this is called decanting), pouring the former upon a filter. Add again 100 cubic centimeters of water to the contents of the dish, boil until well colored, let it settle, and decant upon the filter. Do this four times, which will give 300 cubic centimeters or more of clear liquid. This will prob- 16 ELEMENTARY PRINCIPLES OP CHEMISTRY ably be enough, but if more is needed, it is only neces- sary to add another portion of water and repeat the oper- ation. Try the action of this yellow liquid on the red litmus paper. (d) To show that sulphur is a constituent of this liquid : 'Take about one half test-tubeful, add hydrochloric acid, and heat, holding at the mouth of the tube a piece of filter paper, wet with lead acetate solution. What is the effect ? You also see, reappearing in the liquid, yellow, or nearly white, very finely divided sulphur. The yellow liquid obtained in 18/B may, for present purposes, be considered as a solution in water of lime sul- phide or calcium sulphide, a substance containing lime and sulphur in combination. 18/C 0. The reaction, double exchange, between zinc iodide and lime sulphide. Pour about one half the solution of zinc iodide, prepared in (c), into one of the larger beakers and drop in a piece of red litmus paper. Now, both solutions being hot, pour the lime sulphide slowly into the zinc iodide, a little at a time, constantly stirring, until the paper first turns permanently blue. Be careful to avoid any unneces- sary addition of lime sulphide beyond this point. You see a white insoluble solid, a powder, form in the midst of the liquid, and settle to the bottom. A substance formed in this manner is called a precipitate, and the operation is pre- cipitation. (e) To show that the liquid notv contains lime and iodine in combination, but no zinc nor sulphur : Let the precipitate obtained in settle well to the bottom of the beaker ; then, without disturbing it, pour off as much of the liquid as pos- sible upon a filter. Fill up the beaker with warm water, stir the precipitate, let settle, and again decant the liquid upon the filter, collecting what runs through with the first portion. Label this, filtrate (e), and set it aside. Again fill the beaker with water, stir the precipitate, and set this aside to settle, labeling it, precipitate (e). INTRODUCTION 1? From filtrate (e) take about one half test-tubef ul, and add to this a few drops of lime sulphide solution. The absence of precipitation may be taken as evidence of the absence of zinc, or perhaps, more strictly speaking, of zinc iodide. Take again one half test-tubeful of filtrate (e), and add a 'few drops of lead nitrate solution. The bright-yellow pre- cipitate is recognized as lead iodide, and shows the pres- ence of iodine in combination. Boil the rest of filtrate (e) to dryness in the smaller evap- orating dish, having care to moderate the heat as the liquid thickens, so as to avoid spattering and possibly cracking the dish. When the last of the liquid disappears, the full heat of the flame may be applied. As a result, you see the purple vapor of iodine appear. Continue the heating as long as iodine is liberated, and you have left finally in the dish a white, chalklike powder. That this residue is not zinc oxide, you may see by its color and by adding to the whole, or a portion of it, a very small quantity of water, boiling and testing the liquid with red litmus paper. That this residue does not contain sulphur, you may see by adding hydrochloric acid, boiling, and testing the vapor by paper wet with lead acetate solution. In fact, it is lime, the substance with which you started. (/) To show that the ivhite precipitate (e) contains zinc and sulphur, lut no iodine nor lime : Pour the liquid and the precipitate which was labeled precipitate (e) upon the filter. When the liquid has filtered through, fill the funnel with water, not quite up to the edge of the paper, and when this has passed through fill it a second time. The precipitate is thus washed free of soluble matter. Take out a portion of the wet precipitate and dry it by the heat of the flame, holding it on the spatula blade or the asbestos. Drop some of the dried substance into a test-tube, add hydrochloric acid, boil, and test the vapor as before for hydrogen sulphide. 18 ELEMENTARY PRINCIPLES OF CHEMISTRY Drop another portion of the dried substance into a small evaporating dish, barely moisten it with nitric acid. evaporate the acid by gentle heat, then heat with the full flame. You see sulphur appear and burn, but no iodine, and you have left finally zinc oxide, seen also in Exps. 15/ 3 and 16/2 . Boil another portion of the dried substance with a little water, and test by red litmus to see that it is not lime. EXPLANATORY NOTE. By the operation in 18/A you have combined zinc and iodine, making zinc iodide ; in 18/B you have combined lime and sulphur, forming lime or calcium sulphide ; by mixing the solu- tions of these two substances you have obtained an insoluble solid which contains zinc and sulphur combined that is, zinc sulphide and at the same time you have a liquid which contains lime and iodine com- bined that is, lime or calcium iodide. Thus the two substances, zinc- iodide and lime sulphide, with which you started, have exchanged con- stituents and produced two entirely different substances. Such a change is expressed by the general form : As to Notation. 19 There is used in chemistry, for the purpose of designating substances and describing changes, a system of symbols, which, at this early stage of the study, may best be regarded as a system simply of arbitrary signs, although it really has a much deeper significance. Thus the fol- lowing symbols are used : For sulphur, S ; for oxygen, ; for zinc, Zn ; for iron, Fe (Latin, ferrum); for hydrogen, H ; for sulphur dioxide, S0 a , signifying qualitatively that it is made of sulphur and oxygen, or contains these as sole constituents ; for zinc sulphide, ZnS, and for iron sulphide, FeS, since they contain solely zinc and iron respectively, and sulphur; for hydrochloric acid, HC1, containing hydrogen and chlorine; for hydrogen sulphide or hydrosulphuric acid, H 2 S, since it contains hydrogen and sulphur. The fact that sulphur and oxygen produce by chemical change sulphur dioxide is expressed in equation form thus : S + 2 = S0 2 . The reactions between sulphur and zinc and iron are thus expressed : Zn + S = ZnS. Fe + S = FeS. INTRODUCTION 19 The substance used under the label, hydrochloric acid, is really a 19/1 solution of this substance, HC1, in water. Now, when hydrochloric acid acts on zinc sulphide, the hydrogen of the acid combines with the sulphur of the sulphide forming hydrogen sulphide, and the chlorine of the acid with the zinc of the sulphide forming zinc chloride, ZnCl 2 ; similarly in the case of the iron sulphide. These reactions are ex- pressed by the equations : FeS + 2HC1 = H a S + FeCl 2 . The zinc chloride, ZnCl 2 , and the iron chloride, FeCl 2 , remain dissolved in the water. The numerals which you see used in this system of chem- ical "shorthand" may for the present best be regarded as arbitrary parts of the system. In the " additional illustrations of chemical change " (see Exp. 15/i , etc.) you have used lead, symbol Pb (Latin, plumbum), and iodine, symbol I, and lead iodide, symbol PbI 2 ; mag- nesium, symbol Mg ; zinc, Zn, and oxygen, ; also the oxides of mag- nesium, zinc, and lead, symbols MgO, ZnO, and PbO. What is the meaning of the following equations ? Pb + I 2 = PbI 2 . Mg + = MgO. Zn + = ZnO. Pb' + = PbO. Referring to the experiments under substitution (see Exp. 17/t , 19/2 etc.), you should realize the meaning of this equation, disregarding the numerals : Zn + 2HC1 = ZnCl 2 + H 2 . The symbol for sulphuric acid is H 2 [S0 4 ] ; that is to say, this substance contains hydrogen and something else, the nature of which need not now be considered ; this " something " is represented by the symbol in the brackets. The symbol for zinc sulphate is Zn[S0 4 ], hence the reac- tion between zinc and sulphuric acid is represented thus : Zn + H 2 [S0 4 ] = Zn[S0 4 ] + H 2 . The symbol for nitric acid is H[N0 3 ], and of lead nitrate is Pb[NO s ] 2 ; hence the reaction between lead and nitric acid is represented thus : Pb + 2H[N0 8 ] = Pb[N0 8 ] a + H 2 . Note that these last three equations fall under the general form 20 ELEMENTARY PRINCIPLES OF CHEMISTRY 4. Additional Illustrations of Physical Properties 20 Distillation. See Exp. 24/ 6 , and Appendix, 18, and Part I, tfo. 12. 20/1 Sublimation. Heat gently a small fragment of iodine in a dry test-tube. Describe the phenomenon. Kote the color of the vapor and its odor (cautiously), and its weight compared with that of air (invert the tube). Does the solid melt ? Does the vapor pass through the liquid to the solid form on cooling ? Is the solid which is deposited (the sublimate) crystalline or not ? (See Part I, Xo. 12.) 21 Solution and crystallization. Add a few drops of hydro- chloric acid to a few drops of a solution of lead acetate in a test-tube. What takes place ? (Precipitation.) Boil the contents of the test-tube, adding a little water, if necessary, to dissolve the white powder first formed ; set aside to cool. What takes place on cooling ? EXPLANATORY NOTE. The white powder formed on mixing the two liquids is lead chloride (symbol, PbCl 2 ). Does its formation involve a physical change or a chemical change 1 It is not soluble, or only slightly so, in cold water ; hence it appears as a solid, a powder, in the midst of the liquid. Lead chloride dissolves in hot water, and, as the solution cools, reappears in crystalline form ; this is a purely physical change. 21/4 Solution and crystallization. Dissolve about 10 grams of alum (a porcelain evaporating dish is convenient) in 50 cubic centimeters of hot water. Filter the solution. Dis- solve about twice as much copper sulphate in a similar manner. Filter this. Mix the two solutions; boil down to about one half the volume, and set aside to crystallize (it may be left until the next day). Can you distinguish the two substances in the crystals? Can you distinguish them in the solution ? 21/7 Water of crystallization. Put about 1 gram of copper sulphate crystals in a dry test-tube, and weigh carefully the tube and contents ; to tenths is sufficient. Heat slowly, holding the tube so that the open end is a little the lower. INTRODUCTION 21 What condenses on the cooler portion of the tube ? When the liquid has disappeared let the tube cool, and weigh again. Shake out the white dehydrated substance, and let a drop of water from the finger come in contact with a portion of it. Observe carefully what takes place. Bring a drop or two of alcohol in contact with another portion. The dehydrated, amorphous substance may be redissolved with hot water, and crystallized as it was at first. (See Part I, No. 21/e. ) Heat in similar manner some alum crystals, also some crystals of potassium dichromate, or of potassium nitrate. Efflorescence and deliquescence. Expose to the air for 21/8 twenty-four hours, more or less, (a) some crystals of sodium carbonate or of sodium phosphate ; (b) some calcium chlo- ride or sodium hydroxide (see also Exp. 17/i b). Heat of solution. Determine the effect on temperature 22 (see Appendix, 15) of dissolving in water (a) some com- mon salt that is, sodium chloride (use a beaker half filled with water and about a tablespoonful of salt) ; (b) some sodium hydroxide ; (c) some hydrochloric acid (concen- trated). Melting point, determination of (see Appendix, 16). De- 23 termine the melting point and freezing point of paraffin thus : Prepare a tube of thin wall and small bore by heat- ing in the flame a piece of glass tubing, an inch or two from its open end, until the glass is well softened ; then, draw- ing it out slowly until the bore becomes quite small, and finally applying the flame at the narrow part and separat- ing completely the two portions of the tube. The narrow end is thus closed by fusion. Put into this a few bits of paraffin ; warm just enough to melt the latter-; shake or jar the tube so that the liquid shall completely fill the narrow portion, leaving no air bubble. With a rubber band attach the tube alongside the thermometer so that the paraffin is about opposite the bulb. Suspend the thermometer and * tube over a beaker of water which is supported on an asbes- 19 22 ELEMENTARY PRINCIPLES OF CHEMISTRY tos board or a wire gauze, so that the bulb and paraffin shall be well immersed. Slowly heat the water and observe the temperature when liquefaction is first noted. When this is complete, let cool, and observe the temperature when solidification begins. Repeat four or five times, until the several readings of each point are fairly concordant. As the melting or the freezing point is approached the tem- perature should be allowed to change very slowly, and the water should be well stirred. The freezing point may be somewhat lower than the melting point, due to surfusion. 24 Boiling point, determination of, and conditions affecting. Determine the boiling point of water thus (see Appen- dix, 17) : Use a distilling flask of about 200 or 225 cubic centimeters capacity, with a side delivery tube carrying a short piece of rubber hose. Half fill with water. Through the stopper insert a thermometer (wetting it first) so that the bulb is below the exit and above the surface of the water, high enough to avoid being spattered during boil- ing. Support the whole on asbestos, or gauze, and the iron stand so that there may be no danger of overturning. Heat until the liquid boils and the temperature reaches its maximum, which is steadily maintained for five minutes or more. This is the 'boiling point. It should be, with a good thermometer, 100 C. ; but thermometers are often inac- curate, and may be corrected by actual observation, made with due care. Vary the experiment as follows : 24/1 1. By pinching with the tongs, close the rubber exit tube while the liquid is still boiling, and note the effect on temperature. Let it rise only about 2, then open and ob- serve what takes place. 24/2 2. Lower the thermometer until the bulb is immersed, boil, and note again the maximum temperature, first with the tube open, then with it closed. The temperature of the boiling liquid is likely to be higher than that of the vapor under similar conditions. INTRODUCTION 23 3. Let it cool sufficiently, then add to the water about 24/3 a tablespoonful of salt and a few crystals of copper sulphate, boil, and observe the temperature of the vapor as before (it should be unchanged). 4. Lower the thermometer and take the temperature of 24/4 the boiling solution. Collect some of the liquid (distillate) which drops from 24/5 the exit tube. Does it show any color ? Has it the taste of salt? (Distillation.) Fit a test-tube well with a cork, fill about one third with 24/1 water, and boil. After it has boiled a few seconds and while it is still boiling, remove from the flame, quickly in- sert the cork, invert the tube, immerse the stoppered end just under the surface of some water in a convenient ves- sel, and pour cool water on the upper end. What takes place ? What is the explanation of the phenomenon ? Let cool, uncork, and note the inrush of air. If the cork does not come out easily, force a pin between the cork and the glass and withdraw it. Optional experiment. In a small evaporating dish evapo- 25 rate to dryness some distilled water. Is there any residue ? Similarly evaporate to dryness some sample of natural water. Is there a residue ? CHAPTER II EXPERIMENTS ILLUSTRATING THE FUNDAMENTAL QUAN- TITATIVE LAWS OF CHEMICAL CHANGE 33 NOTE. Study with special care the quantitative relations in the ex- periments of this chapter. In all the measurements of quantity which you make, whether in this work or in the work of subsequent chapters, try to make an estimate of the uncertainty which is necessarily in- volved in the measurements by reason of the conditions in which you work. In some of the problems it will be important to make at least two measurements of the same quantity, in order to show how much uncertainty is thus involved. As the work advances, after you have studied the primary topic which is illustrated, give some thought to the topics already passed, which may find secondary illustration by the experiment in hand, especially as to these general laws, and also as to the kind of reaction involved, and other definitions presented in Chapter I. Have care always to state in your notes the specific topic to be illustrated, and, when practicable, the specific experimental problem to be solved ; then describe your method of solution, endeavoring to give the essential features, apparatus, incidental observations, etc. ; then, as a rule, present the data of observation, carefully labeled, the calculations, if any, and the final result, with your conclusions therefrom. In recording data, be sure that your notes include all original observations. For example, you are to determine the weight of a sub- stance contained in a dish ; you do so by determining the weight of the dish with its contents, and the weight of the dish alone, the differ- ence between these being the weight sought. Now, the important point is to record in your notes, not simply this difference, but the original weights which by subtraction give the desired value. Notes should be written in the laboratory, during the progress of the experiment, and quantitative data should be recorded at once. Do not rush through an experiment and then try to write up your notes from memory, perhaps outside the laboratory. Do not delay recording results until you have learned whether they are good. Put them all 24 QUANTITATIVE LAWS OF CHEMICAL CHANGE 25 down, and later mark " erroneous," if necessary. Do not be discouraged if the result at first trial is unsatisfactory. You will often need to perform an experiment once to learn how to do it, and need to repeat for successful results. It is well to bear in mind that practical chem- istry is more or less a handicraft, ancl that you will surely fail to get good return for your time and labor unless experiments are performed carefully as well as studiously. Make it a rule to read through the directions for the whole experi- ment before you start upon its performance. 1. The Law of Persistence of Mass A. Preliminary. Some apparent contradictions : (a) Place about 1 gram of mercury sulphocyanate in 34/1 & small evaporating dish, counterpoise the dish and its contents on the balance, ignite the substance by a match or a hot iron wire. Describe the phenomenon, and note the effect on the equilibrium of the balance. Had you made the experiment without using the balance, what would you have inferred as to the effect of the change on the quantity of substance ? (b) An experiment for the teacher to perform before the 34/2 class. Burn a taper, or alcohol in a small lamp, in such manner as to collect and weigh the products of combustion. This may be conveniently done as follows : Suspend the taper or very small lamp (which can be made from a short test-tube, a cork, and a piece of wicking) from the lower end of a student-lamp chimney, to the upper end of which is fitted a perforated cork carrying a train of four absorp- tion tubes ; the first two of these are filled with fragments of sodium hydroxide, and the last two with fragments of calcium chloride ; to the farther end of the train is at- tached a rubber hose leading to a filter pump or to an aspirator ; the whole apparatus lamp, chimney, and absorb- ing train is then suspended from the arm of a suitable balance and counterpoised. When this is ready, start the pump which draws a current of air through the apparatus, 26 ELEMENTARY PRINCIPLES OF CHEMISTRY carrying with it the products of combustion into the tubes, where they are retained. Light the lamp and allow it to burn fifteen or twenty minutes, noting the effect on the FIG. 1. Diagram of apparatus for Exp. 34/ 2 . L, alcohol lamp ; C, lamp chimney; TI, T* T a , and T 4 , absorption tubes. equilibrium of the balance. How does the weight of the products of combustion compare with the weight of the material burned? (Fig. 1.) 34/3 An alternative form of apparatus for this experiment may be more simply provided as follows : Make a cylinder of wire gauze which shall fit closely in the upper half of the student-lamp chimney ; fit the lower end with a cork, through which several holes are bored, and to which a short piece of taper may be attached (see Fig. 2), or, preferably, QUANTITATIVE LAWS OF CHEMICAL CHANGE 27 suspend the little alcohol lamp as in Fig. 1. Fill the gauze cylinder with fragments of sodium hydroxide, and suspend the whole by a loop of wire from the arm of the balance. Light the lamp or the taper. Although the products of combustion are not so fully re- tained as in the preceding form, the gain in weight is made evident by burning for a few minutes. EXPLANATORY NOTE. The alcohol and the candle contain carbon and hydrogen as elementary constitu- ents. These in the process of burning combine chem- ically with the oxygen of the air, forming as products gaseous carbon dioxide and hydrogen oxide i. e., wa- ter. These products are carried into the tubes, where the carbon dioxide is retained by the sodium hydrox- ide, and the water is retained partly by condensing to liquid form, and partly by the strongly hygroscopic substance, calcium chloride. The products of the combustien, therefore, weigh more than the material burned by just the weight of the oxygen taken into combination from the air. All the ordinary fuels likewise contain carbon and hydrogen ; therefore the products of ordinary combus- tion are carbon dioxide and water, and combustion itself is simply an instance of chemical action accompanied by heat and light. (Compare with Exps. 13, 15/ 2 and 15/ 3 .) B. To illustrate the law: (1) Eecall or repeat Exps. 13/ 2a , and 15/ lt (2) In a gas generator (see Appendix, 6) fitted with a 34/5 thistle-tube and a delivery tube place a charge of marble (i. e., calcium carbonate, symbol CaC0 3 ) in small lumps, with enough water to seal the end of the thistle-tube ; in another vessel (a bottle or a deep beaker) place a strong solution of sodium hydroxide ; in a third (a small beaker), some concentrated hydrochloric acid and a glass rod. Counterpoise the three vessels with their contents on the heavier balance (see Appendix, 1, A). Cautiously turn the acid, a few drops at a time, into the generator (observe 28 ELEMENTARY PRINCIPLES OF CHEMISTRY what takes place), letting the reaction go on slowly through a measured interval of time while the delivery tube opens into the air, and until the effect on the equilibrium is un- mistakable. (What is the effect ?) Then drop the delivery tube into the second vessel so that the gas bubbles through the solution ; again counterpoise, and continue the reac- tion at about the same rate, and during the same length of time as before ; then again note the effect on the equi- librium. Secondary observation. To show that an invisible gas, quite different from air and not before present, is produced by this change, remove the material from the balance, drop the delivery tube into a clean beaker or bottle, partly covered by a glass plate (see Appendix, 19, V), and pour a little more acid into the generator. After the effervescence has continued a few seconds, slide the plate to one side and plunge a lighted match into the bottle. The gas produced is carbon dioxide, symbol C0 2 . EXPLANATORY NOTE. This experiment, owing to the limitations in weighing, can give but crude quantitative results, showing simply that the balance, which gives no indication of differences less than 0.2 or 0.3 of a gram, indicates unmistakable loss of weight in the first condi- tions, but no loss in the second, although the reaction goes on as in the first, The chemical name of marble is calcium carbonate, symbol CaC0 8 . It is a salt derived from the metal, calcium, and carbonic acid, symbol H 2 C0 3 . By the action of hydrochloric acid on this substance, calcium chloride, another salt, symbol CaCl a , is produced, and carbonic acid. The latter substance breaks up at once into carbon dioxide, symbol C0 3 , and water, symbol H 2 0. The carbon dioxide passes into the atmosphere as an invisible gas, the water adds itself to the rest of the water, and the calcium chloride remains in solution. These two reac- tions are expressed in equation form as follows : (1) CaC0 3 + 2HC1 = CaCl a + H a CO. (2) HaCOs = C0 3 + H 2 0. Sodium hydroxide, symbol NaOH, is a base (see Part I, 30/ a ), and when the carbon dioxide comes in contact with its solution in the sec- QUANTITATIVE LAWS OF CHEMICAL CHANGE 29 ond bottle the two substances combine and another salt is produced, namely, sodium carbonate, symbol Na 2 C0 3 . This remains in solution. The following equation expresses this reaction : SNaOH + C0 2 = Na 3 CO s + H 2 0. 2. The Law of Fixed or Definite Proportions (a) Evaporate to dryness a few drops of hydrochloric 37 acid solution in the small evaporating dish, heating gently with the caution necessary in this operation (see Appendix 12). Acid of a good degree of purity is needed. (b) Likewise evaporate a few drops of ammonium hy- droxide solution. Describe the results. (1) Carefully weigh to tenths of a gram the small porcelain dish with a short glass rod, both articles being clean and dry. [It will be advantageous to have two or even three weighed dishes to use. They may be distinguished by scratches on their edges made with a file.] With some suitable measure (a test-tube with a slip of gummed paper or a file scratch on its side will serve, although a burette or graduated pipette is better) take two equal volumes of the ammonium hydroxide solution, pour- ing them into a weighed dish. Exactly neutralize this liquid with the hydrochloric acid solution, using litmus paper as indicator, and adding the acid drop by drop, stir- ring after each addition, until the paper just turns per- manently pink. Note carefully the volume of acid used. Evaporate the contents of the dish to dryness on the water- bath ; this is facilitated by frequent stirring toward the end of the operation. Let cool, and determine the weight of the salt obtained. How will you make sure that the salt is completely dried ? Redissolve the salt, using very little water; filter if neces- sary, and crystallize. Dry the crystals on filter paper, and test them with blue and with red litmus paper slightly wet. 30 ELEMENTARY PRINCIPLES OF CHEMISTRY (2) Take the same volume of hydrochloric acid which was used in (1), add to it only one measureful of ammonium hydroxide, evaporate, and weigh as before the salt obtained. Eedissolve, filter, and crystallize this salt. Test the crystals with red and with blue litmus paper. (3) Take again the same volume of acid as previously used, add to it three measurefuls of ammonium hydroxide, evaporate, and weigh as before. Eedissolve, filter, and crystallize this salt. Test the crystals with red and with blue litmus paper. 38 QUESTIONS. Are the three portions of salt thus obtained different samples of the same substance, or are they different substances? What is the ratio between the quantities of salt obtained in the three cases ? Have you any evidence that hydrochloric acid or ammonium hydroxide passes off during the evaporation, in (1) ? in (2) ? in (3) ? What is the logic of the experiment ? Reason it out fully. Is heat liberated when hydrochloric acid and ammonium hydroxide are mixed ? 39 EXPLANATORY NOTE. The salt obtained by the combination of hydrochloric acid and ammonium hydroxide (symbol NH 4 OH, a base) is named ammonium chloride, symbol NH 4 C1, and the reaction is thus expressed : Hydrochloric acid and ammonium hydroxide produce ammonium chloride and water ; or in equation form, HC1 + NH 4 OH'= NH 4 C1 + H 2 O. 3. The Law of Multiple Proportions 40/a (a) Take of iodine and of mercury in the ratio of 252 parts by weight of the former and 199 of the latter ; for the actual experiment weigh out accurately (in some glass vessel, as both substances attack metals) 6.30 grams of iodine, and 5.00 grams of mercury. Transfer the mer- cury to a mortar (previously weighed with its pestle on the heavier balance). Add a few drops of alcohol, then a small portion of the iodine, and rub gently with the pes- tle ; then add another portion of iodine, and rub ; and so on until all the iodine is used and the whole is thoroughly mixed. QUANTITATIVE LAWS OF CHEMICAL CHANGE 31 It is well to keep the mixture slightly moist with alcohol during only the first stages of the operation, in order to avoid overheating, otherwise there may occur a slight flash in the mortar, accompanied by the fusion of the substance. If this happens, it is better to clean the mortar (use sand to scour), and start the experiment anew. Avoid in- haling the fumes of iodine, and protect the hand, if desired, with a fold of the towel. Iodine stains may be removed with sodium sulphite or ammonium sulphide. Persistent rubbing may be necessary even after the sub- stances seem well mixed. The operation when completed should yield a bright-red powder, mercuric iodide, symbol HgI 2 . Weigh the mortar and contents on the heavier balance. Take out a small sample and apply alcohol to it in a 40/1 test-tube. [The alcohol should not be colored by the iodine ; if it is so colored, continue the rubbing.] Warm the alcohol to gentle boiling, having care that the alcohol vapor is not ignited by the flame ; set aside to cool. The amorphous powder dissolves in hot alcohol, and deposits on cooling in crystalline scales, sometimes bright scarlet, sometimes yellow (allotropic forms). Sublime another small portion of the red powder in a 40/2 dry test-tube, note the red and yellow sublimate (allotropic forms) ; rub the yellow with a glass rod. Note the crystal- line residue after fusion. (V) Take 5.60 grams of mercuric iodide obtained in 40/b (), which is equal to ^ (6.3 + 5) and 2.50 grams of mer- cury. Add the mercury in small portions at a time to the mercuric iodide in a mortar, rubbing thoroughly after each addition. This should yield a greenish-yellow powder (mercurous iodide, symbol Hgl). If the color is not satis- factory, continue the rubbing and, perhaps, allow to stand for twenty-four hours. Weigh the mortar and contents on the heavier bal- ance. Treat a small portion as in (a) with boiling alcohol. 40/3 Does it dissolve ? 32 ELEMENTARY PRINCIPLES OF CHEMISTRY 40/4 Sublime another portion ; note the character of the sub- limate and the residue after heating. 40/5 EXPLANATORY NOTE. Mercurous iodide is insoluble in hot alcohol and decomposes on heating into mercuric iodide and mercury. QUESTIONS. How do you differentiate between mercury and iodine mixed, and mercuric iodide! between mercuric iodide and mercurous iodide f between mercuric iodide and mercury mixed, and mercurous iodide f How does the weight of the mercuric iodide obtained com- pare with that of the mercury and iodine used 1 What becomes of the alcohol f How does the weight of the mercurous iodide compare with that of the mercuric iodide and the mercury used ? What is the ratio between the quantities of mercury in the two substances reckoned for constant quantity of iodine ? between the quantities of iodine, reckoned for constant quantity of mercury ? What is the logic of the experi- ment f Show clearly that it is not contradictory to Law 2. What does Law 2 affirm in regard to these two reactions ? What does Law 1 affirm with regard to them ? Is heat liberated in the reaction be- tween mercury and iodine f 4. The Law of Equivalent Proportions* 41 General problem. To investigate the relation between those quantities of different substances which produce equal chemical -effects. Specific illustrative problem. A. To determine the rela- tion between the quantities of oxygen which combine respectively (1) with 2.40 grams of magnesium, and (2) with 6.50 grams of zinc. B. To determine the relation between the quantities of hydrogen liberated from hydrochloric acid respectively (3) by 2.40 grams of magnesium, and (4) by 6.50 grams of zinc. * Inasmuch as the experiment under this topic involves the meas- urement of gas- volume, the instructor may prefer to introduce the laws of Boyle and of Charles, Chapter IV, at this point. The writer prefers to give them simply as arbitrary rules and to study them later, rather than to interrupt the logical development of this chapter. QUANTITATIVE LAWS OF CHEMICAL CHANGE 33 A 1 To determine the quantity of oxygen which combines 41/1 with 2.40 grams of magnesium : Recall Exp. 15/ 2 . Weigh to hundredths of a gram a porcelain crucible with its lid, clean and dry. Weigh out exactly 1.00 gram of magnesium ribbon. Breaking this into small pieces, transfer the whole to the crucible. Support the latter, covered by its lid, on the pipe-stem triangle and iron stand (see Appendix, 11). Apply heat, slowly at first, then the full heat of the flame. Raise the lid slightly from time to time to see what is taking place and to let air in, but avoid losing any of the white smoke. When the danger of this is passed, remove the lid and continue the heating about fifteen minutes. When the crucible is cool, weigh it with its lid and con- tents to hundredths. To make sure that the reaction is complete, heat, cool, and weigh again. When heated, the magnesium combines with the oxygen of the air and the product is magnesium oxide, MgO. Calculate from the mass of oxygen which combines with 1.00 gram of mag- nesium what mass combines with 2.40 grams of the same. (Upon what law is this calculation based ?) At least two determinations of this value should be made. Secondary items. Describe the phenomenon seen and 41/a the substance produced. Does the reaction evolve heat ? Does magnesium oxide dissolve in hot water ? Test with red litmus paper. Does it dissolve in dilute sulphuric acid ? Boil it in the crucible or dish with a very little of the dilute acid, filter, and crystallize the salt, magnesium sulphate, MgS0 4 (compare Exp. 17/ 2 ). How does the reaction be- tween magnesium oxide and sulphuric acid, symbol H 2 S0 4 , differ from that between magnesium and the same acid ? An Alternative Method This method involves making, first, magnesium iodide, Mgl a , then 41/b heating this substance, by which it is decomposed into iodine and mag- 34 ELEMENTARY PRINCIPLES OF CHEMISTRY nesium, and the latter is converted into the oxide, MgO, all without loss of substance other than the iodine. The final result is the same as in 41/i, but it is brought about with less elevation of temperature. Weigh to hundredths a small evaporating dish with a short glass rod. Weigh out carefully 0.50 of a gram of magnesium, and place this in the dish. Weigh out about 5 grams of iodine. Pour on the magnesium enough alco- hol to cover it. Then add the iodine, a little at a time, stirring constantly. With a little care the reaction may be so controlled that the heat evolved shall not cause spatter- ing and consequent loss of material. When the magnesium is entirely acted upon, heat the dish and contents on the water-bath until only a thick, siruplike mixture is left. Then transfer the dish to the iron ring and heat with a very small flame, holding the burner in the hand (as before directed for careful evaporation ; see Appendix, 12), and stirring constantly. Continue this until the danger of spattering is past, then apply the full heat of the flame. As this causes iodine fumes, it may be necessary to perform this part of the experiment in the hood. Continue the heating for a few minutes after the iodine has entirely dis- appeared. Finally let cool, and weigh the dish and con- tent, which is magnesium oxide, a fine powder only slightly brown in tint. The calculations are the same as in the first method. A2 41/2 To determine the quantity of oxygen which combines with 6.50 grams of zinc : It is impracticable to do this by direct combustion, as in the mag- nesium experiment. The method is to act on zinc with nitric acid, pro- duce zinc nitrate, dry this, then ignite it and thus produce zinc oxide without loss of other than volatile material. Refer to Exps. 15/ 3 and 16/ 2 . Weigh a small evaporating dish with a short glass rod to hundredths of a gram. Weigh out 2.00 grams of granu- lated zinc. Transfer the latter to the dish. Pour a drop QUANTITATIVE LAWS OF CHEMICAL CHANGE 35 or two of nitric acid down the rod, letting it come slowly in contact with the zinc. (Remember that nitric acid is very corrosive.) When the first violent reaction is over, add a few drops more of acid, and so on until the zinc is completely dissolved with the least acid that will serve. Evaporate the contents of the dish to dryness on the water- bath. The evaporation may be performed more quickly, but only with closer attention, by holding the burner in the hand, having the flame low, and applying just enough heat to keep the liquid gently boiling, until there is danger of spattering, and then not enough heat to cause the for- mation of bubbles (see Appendix, 12). When the salt is thoroughly dry, increase the heat, and finally give it the full flame for about fifteen minutes. When the reaction is finished, let the dish cool, and weigh it with its contents to hundredths of a gram. The product is zinc oxide. To make sure that the reaction is complete, heat, cool, and weigh again. From the mass of oxygen which combines with 2.00 grams of zinc, calculate what mass combines with 6.50 grams of the same. At least two determinations of this value should be made. Secondary items, Describe the reaction between zinc and 41 /c nitric acid. Is heat evolved ? What takes place when zinc nitrate is heated ? Does zinc oxide dissolve in water ? in dilute sulphuric acid ? Optional experiment : Filter, and crystallize the zinc sulphate (see Exp. 17/ 8 ). How does the reaction between zinc oxide and sulphuric acid differ from that between zinc and the same acid? What is obtained by the action of hydrochloric acid on zinc oxide ? Summary. Compare with each other the individual 41/d quantitative results with magnesium, also those with zinc ; then compare the average of the first with the average of the second. What is the conclusion as to the masses of oxygen which combine respectively with 2.40 grams of magnesium and with 6.50 grams of zinc ? 36 ELEMENTARY PRINCIPLES OF CHEMISTRY B 3 41/3 To measure .by volume and calculate into weight the quantity of hydrogen liberated from hydrochloric acid by 2.40 grams of magnesium : (a) To measure the capacity of the collecting bottle, No. x (a bottle with glass stopper holding about 2,500 c.c.): Fill the tank with water, which should be brought to the temperature of the room, if necessary by heating portions in the iron vessel. Fill the collecting bottle completely and put the stopper in place. Measure the volume of this water in cubic centimeters ; this may be done by pouring, without loss, into a graduated cylinder, or, if such vessel is lacking, as follows : Fill a graduated flask repeatedly from the bottle (one holding 750 c.c. is convenient), and weigh, in a vessel not unnecessarily large or heavy, the last residue which is insufficient again to fill the flask ; the weight of the water in grams may be taken with only inconsiderable error as its volume in cubic centimeters. Determine thus the capacity at least twice. The variations should not be more than one or two cubic centimeters. [Why not weigh at once the bottle filled with water ?] 41/3b (b) To generate the gas : Weigh out carefully 2.40 grams of magnesium ribbon, transfer without loss to the gas generator (see Appendix, 6). Add water just sufficient to seal the thistle-tube ; test for leakage by blowing in the exit end of the delivery tube until the water rises four or five inches in the thistle-tube, then pinching the rubber hose, or closing the end of the tube with the tongue. The column of water should maintain its height, or sink only slowly. A film of water between stopper and glass will help to prevent leakage, or, as a last resort, the stopper may be sealed with vaseline or with paraffin. Have some strong hydrochloric acid in a beaker. It will be found convenient to let a test-tube lie in the bulb of the thistle-tube ; it may be used like a rod to aid in QUANTITATIVE LAWS OF CHEMICAL CHANGE 3? pouring from the beaker, and the rounded end makes a kind of valve. (c) To collect and measure the gas : Fill the collecting 41/3c bottle with water, put in the stopper, invert in the tank, remove the stopper, incline the bottle so that it rests on the wall of the tank and slip into its mouth the bent delivery tube (compare Appendix, 19, I). Pour acid, a little at a time, into the thistle-tube until no more gas is generated. As much heat is liberated in this reaction, it is well to set the generator down into the water of the tank. Ascertain now or later the volume in cubic centimeters of the acid turned in. When the effervescence ceases, raise the bottle so that the water within and without is as nearly as practicable at the same level, withdraw the delivery tube, insert the stop- per under water, and remove the bottle from the tank. Measure the volume of the water left in the bottle, and calculate the volume of the gas collected. Observe the temperature of the water in the tank, which should be that of the room, so that it may be safely assumed that the temperature of the gas is the same. Observe by the barometer the pressure of the atmos- phere (see Appendix, No. 15 B). Take note that hydrogen with air may make a danger- ously explosive mixture ; therefore have great care that the hydrogen is not brought near a flame. Calculations. (See Appendix, 20.) The data of this 41/8d problem are : The weight of metal used ; The total capacity of the bottle ; The volume of acid turned in (how is this used ?) ; The volume of water left in the bottle ; The temperature and pressure of the gas. From these is first calculated the volume of the gas generated, measured in cubic centimeters at the observed temperature and pressure. 20 38 ELEMENTARY PRINCIPLES OF CHEMISTRY Now, since change of temperature and change of pres- sure cause change in the volume of gases, it is important that observed volumes, in order to be comparable, should be reduced to standard conditions of temperature and pres- sure. It is customary to use as standards C. for tem- perature and 760 millimeters of mercury column in the ba- rometer for pressure. The method of reducing observations to the standards may be taken arbitrarily now, to be ex- plained in Chapter IV. The observed volume is reduced to volume at by applying the law that " a volume of gas at increases by -^ of itself for each degree of rise in temperature, pressure remaining unchanged." The observed volume is reduced to volume at 760 mm. by multiplying by the observed pressure in millimeters and dividing by 760 ; but since the gas is collected and measured over water, it is saturated with water vapor, and the observed pressure is made up of the true pressure of the gas plus the pressure of the vapor at the given temperature. The pressure or tension of the water vapor at different temperatures is obtained by observation, and given in tables accessible in books of reference ; this value for the ordinary temperature of 20 is 17 mm., and may be subtracted from the observed pressure as a slight correction. (See Appendix 22.) The volume of gas generated, thus brought to standard conditions, is finally calculated into weight, by aid of the fact that one liter of hydrogen at and 760 mm. weighs 0.0899 of a gram. Does the fact that the hydrogen collected in the bottle is mixed with air coming from the generator affect the problem ? At least two fairly concordant determinations should be made of the hydrogen by magnesium. B 4 41/4 To measure by volume and calculate into weight the quantity of hydrogen liberated by 6.50 grams of zinc (com- pare Exp. 17/j) : Proceed as in B 3, () and (c), using 6.50 QUANTITATIVE LAWS OF CHEMICAL CHANGE 39 grams of granulated zinc. At least two fairly concordant determinations of this value should he made. Summary. Compare with each other the individual 41/5 results from magnesium ; also those from zinc ; then the average of the first with the average of the second. What is the conclusion as to the quantities of hydrogen liberated respectively by 2.40 grams of magnesium and by 6.50 grams of zinc ? Compare the results in A and B. Are the masses of magnesium and zinc which produce equal chemical effect in combining with oxygen also the masses which produce equal effect in liberating hydrogen ? 5. The Law of Gas-volumetric Proportions General problem. To determine the proportions by vol- ume in which gaseous substances react. Specific illustrative problem. To investigate the propor- 47 tion by volume in which nitrogen dioxide reacts (1) with air and (2) with oxygen. EXPLANATORY NOTE. The plan of work under this topic must be to state first the fundamental fact involved, and to let the experiment serve rather to illustrate the practical application of this fact to the observations in hand. Nitrogen dioxide, symbol NO (named also nitric oxide), combines with oxygen on contact in the ratio of 4 volumes (gaseous) of the former to 2 volumes (gaseous) of the latter, forming 2 volumes (gaseous) of nitrogen tetroxide, symbol N a 04, a substance which is soluble in water. Air is a mixture of oxygen and nitrogen, 1 volume of the for- mer to 4 volumes of the latter. Nitrogen does not react with nitrogen dioxide. (a) To calibrate two gasometric (i. e., gas-measuring) 47/a tubes : Use two test-tubes, 230 mm. (9 in.) long and 18 mm. (f in.) wide, and as a unit tube one about 44 mm. (If in.) long and 10 mm. (f in.) wide. Divide the large tubes into portions of equal capacity by filling the small tube with water, pouring its content into the large tube, and marking the level of the water with a rubber band. 40 ELEMENTARY PRINCIPLES OF CHEMISTRY 47/b (#) To prepare nitrogen dioxide : Put into the generator a charge (to about half fill) of ferrous sulphate (commer- cially known also as copperas and as green vitriol) in small lumps, with water just enough to seal the thistle-tube. Have strong nitric acid with a rod in a small beaker. Turn the acid a little at a time into the generator. Collect the gas over water in the small collecting bottles (see Appendix, 19, I), rejecting, as probably impure, the first three or four bottlefuls ; then retain for use stored in the bottles. Care- fully avoid inhaling the gas; also handle the acid with care, as it is very corrosive. 47/1 (1) By means of the gasometric tubes, measure out vol- umes of air and of nitrogen dioxide in the ratio of 5:2; pouring the latter under the water of the tank, upward from the bottle into the tube (see Appendix, 19, II). It is advantageous to use large rather than small quantities. Pass the content of one tube, under water, into the other (what takes place?), and, after sufficient interval, read the volume of the residual gas. Eepeat the operation enough to get fairly concordant results. If the gas is pure and the measurements well made, the result is 4 volumes in proportion to the 5 and the 2 which were taken ; there- fore 3 volumes, or f of the mixture, disappear in the reac- tion. (What is the cause of this contraction ?) Take a measured sample of the residual gas, add to this a measured volume of air, and observe the contraction ; if it is not zero, make repeated additions, until there is no further contraction. Take another measured sample of the residual gas, add to it a measured volume of nitrogen dioxide ; if the con- traction is not. zero, continue the additions until it is. Does the first residual gas contain oxygen ? Does it con- tain nitrogen dioxide ? What is it ? 47/2 (2) Measure out, as in (1), volumes of oxygen and of nitrogen dioxide in the ratio of 1 : 2. Mix them and meas- ure the residual volume and the contraction. QUANTITATIVE LAWS OF CHEMICAL CHANGE 41 If there is a residual gas, test it for further contraction by repeated additions of oxygen. What is indicated here by contraction or absence of contraction ? Likewise test the residual gas for further contraction by additions of nitrogen dioxide. What does contraction signify here? Why are observations of temperature and pressure unnecessary? (For a convenient form of record see Prob. 8, No. 65/ 6 , Part I.) APPLICATIONS. Measurement of contraction, in conditions like those 47/3 just indicated, may be used to determine the real volume of oxygen or of nitrogen dioxide in a given sample that is, the purity of the sample provided only that there is nothing present to cause contraction in the given conditions, except these two substances. It is only necessary to make sure that the substance to be measured has been entirely used in making the observed contractions. (How is this made sure f) Then the volume of oxygen actually present may be obtained by the propor- tion 3 : 1 :: observed contraction : volume of oxygen (why?); and the volume of nitrogen dioxide may be obtained by the proportion 3:2:: observed contraction : volume of nitrogen dioxide (why ?). Apply this to determine the percentage purity of your sample of nitrogen dioxide and of oxygen. Your sample of the first will be of good quality if you reject enough to have it free from nitrogen coming from the air. Com- mercial oxygen is likely to be impure, also, from the presence of ni- trogen. EXPLANATORY NOTE. The reaction by which the nitrogen dioxide 47/4 is produced is rather complicated for complete explanation at this stage ; the following may suffice : The proximate (see No. 26, Part I) constituents of nitric acid are water and nitrogen pentoxide ; the latter very readily gives up a portion of its oxygen to other substances, and becomes a lower oxide of nitrogen (example of multiple proportions), in this instance the dioxide ; the agent which takes this oxygen from the nitrogen pentoxide is ferrous oxide, a proximate constituent of the ferrous sulphate ; in thus taking oxygen this substance is converted into a higher oxide of iron namely, ferric oxide (example of multiple proportions). Substances which thus give up the whole or a part of their oxygen 48 are called oxidizing agents ; those which thus take on oxygen are called reducing agents. 42 ELEMENTARY PRINCIPLES OF CHEMISTRY 6. The Law of Persistence or Conservation of Energy applied to Chemical Phenomena Heat Disturbance in Chemical Reactions (The student should read Nos. 50-55, Part I, before undertaking the experiment.) 50/1 THE UNIT OF HEAT. The unit of quantity which is used in meas- uring heat is defined as the quantity of heat which raises 1 gram of water 1 in temperature, or, strictly defined, from to 1 C., but, as more commonly used, from 18 to 19. This unit is named the calorie (or gram-calorie). For convenience, a unit equal to 1,000 calories is often used, the kilogram-calorie, designated by the abbre- viation, Cal., while for the smaller unit the abbreviation, cal., is used. 50/2 The quantity of heat disturbance is most simply measured when the reaction takes place quickly and in water solution, as in the neutrali- zation of bases and acids. The result is commonly reckoned for the combining weights in grams of the substances used. (Recall in what reactions you have noted the liberation of heat.) 50/3 Illustrative problem. To determine in calories the quan- tity of heat liberated in neutralizing 125 grams (its com- bining weight) of oxalic acid with ammonium hydroxide, both being in water solution. Weigh out 5.00 grams of oxalic acid ; dissolve this in 100 c.c. of water in a beaker. Take sufficient ammo- nium hydroxide solution to neutralize this acid,* and add enough water to bring its volume up to 100 c. c. Have the two solutions in beakers, and at a temperature not far dif- ferent from that of the room. Observe carefully the tem- perature of each solution. Quickly pour one into the other, stir with the thermometer, and observe the resultant temperature. (The salt, ammonium oxalate, may be read- ily crystallized from this solution by. slight concentration.) * The quantity will depend on the strength of the solution. It is recommended that the instructor determine this for the class, or assign it to some student as a problem. QUANTITATIVE LAWS OF CHEMICAL CHANGE 43 Calculations. In dilute solutions it may be assumed, for 50/4 calculations such as these, that the quantity of heat needed to raise the solution one degree is the same as that needed to raise the water which it contains one degree. Sometimes the weight of the water is taken, and sometimes the volume of the solution is reckoned as if it were pure water. The latter method is somewhat the simpler. In this experi- ment, the two volumes being equal, the temperature of the mixture 'would be the mean of the temperatures before mixing, if no heat were liberated. The difference between this mean and the observed resultant temperature is the number of degrees that the 200 c.c., reckoned as so many cubic centimeters or grams of pure water, have been raised by the heat of the reaction. Calculate from this the cal- ories of heat which would be liberated by neutralizing 125 grams of oxalic acid in dilute solution. Some heat is of course taken up by the material of the containing vessel ; how much, it would be necessary to determine, if greater accuracy were demanded. At least two determinations should be made of this heat of neutralization. CHAPTEE III COMBINING WEIGHTS-NOTATION-EQUATIONS- STOICHIOMETRY NOMENCLATURE No experiments. 61 to CHAPTEE IV EXPERIMENTS ILLUSTRATING THE RELATION BETWEEN VOLUME, PRESSURE, AND TEMPERATURE OF GASES 1. The Law of Boyle Relation between Volume and Pressure of Gases 66 Specific problem. To investigate the relation between the volume and the pressure of a confined mass of air, its temperature remaining constant. Use a gasometric tube of 25 cubic centi- meters in capacity, graduated in fifths (the inside must be dry), filled about one third or one half with air, the rest with mercury, and inserted, the open end down, in a narrow cyl- inder containing mercury. Hold the tube, using a paper holder to avoid changing the temperature by contact with the hand, in a fixed position ; observe the volume occupied by the gas, reading from the top of the mer- cury column ; measure also, the position be- ing unchanged, the length of the mercury column from its top to the free surface of the mercury in the cylinder, using a metric rule or a common one, according to convenience. Now change the position of the tube verti- cally, read the new volume, and measure the new length of column. In this manner make the observa- tions in five or six different positions of the tube. 44 FIG. 3. Appa- ratus to show Boyle's law. VOLUME, PRESSURE, AND TEMPERATURE OP GASES 45 Calculations. It will be found convenient to record the 66/1 volumes in a vertical column, and the lengths in a column parallel with this, each corresponding horizontally with its proper volume. Deduct the several lengths of mercury column from the length of the mercury column in the barometer, and the differences thus obtained measure the relative pressures of the gas when it occupies the corre- sponding volumes. Xow, looking for a relation between the numbers measuring volume and those measuring pres- sure, multiply each volume by its corresponding pressure and compare the products. (See Appendix, 15 B.) How comes it that pressure is measured in linear units ? And why is the observed pressure obtained by subtracting the observed length of column from the barometric length ? What is the limit of accuracy in your measurement of volume Of pressure? How much variation from constancy may the products show and still prove constancy within the limits of observational accu- racy? What does constancy in the products prove concerning the factors ? 2. The Law of Charles Relation between Volume and Temperature of Gases Specific problem. To investigate the relation between 67 the volume and the temperature of a confined mass of air, its pressure remaining constant. Use an apparatus of the following description : A gaso- metric tube of 25 cubic centimeters in capacity, graduated in fifths (the inside must be dry), filled about one third or one half with air, and the rest with mercury. This is in- verted, and the open end inserted in a shallow vessel con- taining mercury. Outside the gasometric tube is a larger glass tube, serving as a jacket, which is closed at the lower end by a stopper, through which pass centrally the gaso- metric tube and, either side of this, one small tube to carry steam and one to drain off the water. A thermometer is suspended beside the gasometric tube from a stopper which ELEMENTARY PRINCIPLES OF CHEMISTRY loosely fits the upper end of the jacket. A flask is con- nected so that steam may be passed into the jacket at will, and the whole apparatus is fixed firmly in position. When it has stood long enough to make sure that the inclosed air has the temperature indicated by the thermometer, read the volume, pressure, and tem- perature. Then apply heat to the flask containing wa- ter, and pass steam into the jacket until the thermom- eter reaches its maximum and holds it steadily for fifteen or twenty minutes. Finally, read again the vol- ume, pressure, and temper- ature. Calculate the aver- age increase in volume for 1, on unit volume, at the initial temperature, pres- sure being constant. To put this result into conventional form, so that it shall be comparable with the result as given in books, it is only necessary to assume the same rate of change between the initial temperature and 0, that you have observed between the initial and final tem- peratures, and from this, to calculate the average increase in volume per degree on unit volume at 0. This is called the increment of volume, also the coefficient of expansion. [What would be the effect of moisture in the gasometric tube on your experimental result ?] FIG. 4. Apparatus to show Charles's law. A, jacket ; B, gasometric tube containing air ; T, thermom- eter ; 8, rubber tube to convey steam from the flask, F ; W, tube to carry off water. CHAPTEE Y EXPERIMENTS ILLUSTRATING THE RELATION BETWEEN EQUIVALENT AND COMBINING WEIGHTS AND CERTAIN SPECIFIC PROPERTIES 1. The Law of Gay-Lussac Relation between Equivalent and Combining Weights of Gases, Elementary and Compound, and their Specific Gravities Specific problem. To determine experimentally the spe- 71 cine gravity (A) of oxygen, (B) of carbon dioxide, and to investigate the numerical relation between these values and the combining weights of the substances, assuming that the formula of the latter is C0 2 . [What is the specific gravity by definition ? (see No. 7, Part I.) What are these combining weights, numeric- ally?] Assume that the weight of one liter of hydrogen at and 760 millimeters is known to be 0.0899 of a gram. The experimental problem then becomes, to determine the weight of one liter of the gas in question at and 760 millimeters; or, if preferred, to determine the weight of any measured volume of the gas at an observed tempera- ture and pressure. The most direct method of solution would be to weigh a mass of the gas in a vessel of meas- ured capacity ; but the experimental difficulties in doing this make it impracticable for beginners. (Can you sug- gest some of the difficulties?) A less direct method is therefore followed. 47 48 ELEMENTARY PRINCIPLES OP CHEMISTRY To determine the weight of one liter of oxygen : 71/A Use a large test-tube (about 230 millimeters or 9 inches long) fitted with a rubber stopper, carrying a glass and rubber delivery tube, like that of the gas generator. The test-tube must be thoroughly dry. Weigh out roughly about 5 grams of manganese dioxide (symbol Mn0 2 ). Place this in the porcelain crucible and ignite it with the full heat of the flame for ten or fifteen minutes to insure complete dryness. Weigh out roughly about 20 grams of powdered potas- sium chlorate (symbol KC10 3 ). Mix the two substances thoroughly in a clean mortar ; transfer to the test-tube, and wipe any adhering dust from the upper part of the tube. Insert, to lie just back of the stopper, a loose plug about 15 millimeters (one half inch) wide, made of a coiled strip of asbestos board or of glass wool. Weigh carefully, to hundredths, the tube and contents thus prepared. Support the tube horizontally between two rings of the iron stand, and tap it gently so that the powder shall lie somewhat spread out. Insert the stopper and connect with the large collecting bottle in the tank. Eecall the experiment (No. 41/ 3 c) with hydrogen under Law 4, Chapter II, and follow here the details of manipu- lation there directed for the collection and measurement of the gas generated. Heat the mixture gently so as to avoid too rapid evolu- tion of gas. It is best to apply heat to only a small por- tion of the mixture at a time, for, if the whole mass be- comes heated, the evolution of gas will not cease when the flame is withdrawn, and more than the bottleful may be generated. When nearly sufficient .gas to fill the bottle has been collected, withdraw the heat, and when bubbles cease to pass, remove the stopper, lest water be drawn into the COMBINING WEIGHTS AND SPECIFIC PROPERTIES 49 tube ; let the tube cool, and carefully weigh it. The loss in weight is taken as the weight of the oxygen which has passed out of the tube. The volume of this gas is measured in the collecting bottle, as before indicated. It is well to shake the gas with the water remaining in the bottle, and then to reopen the bottle for a second under water. (Why?) Calculations, From this weight and this volume calcu- 71/1 late the weight of one liter of oxygen at and 760 milli- meters, and from this value calculate the specific gravity of the gas and compare it with its combining weight. Make a second determination of the same. The origi- nal charge will be more than sufficient to generate two bottlefuls. Secondary observation. After the second determination, 71/2 collect some of the gas in a small bottle and test it with a lighted match. (See Appendix, 19, I and IV.) EXPLANATORY NOTE The potassium chlorate, symbol KC10 3 , by heating loses its oxygen and becomes potassium chloride, symbol KC1. The temperature at which this decomposition takes place is lowered by the presence of the manganese dioxide, hence its use. B. First Step To determine the weight of the carbon dioxide gener- ated from a weighed quantity of calcium carbonate by the action of an acid : Weigh out with care 5.00 grams of this substance in pow- 81/1 der. Transfer it without loss to a filter paper, wiping off any dust adhering to the glass crystal with a small piece of paper. Wrap the whole in a small package, and put it in a beaker. Weigh this beaker with its contents. In a second beaker put about 20 grams (15 c. c.) of nitric acid and a glass rod, and weigh this beaker with its contents. Cautiously pour acid from the second beaker into the first, a little at a time, until there is no further action. Then weigh 50 ELEMENTARY PRINCIPLES OF CHEMISTRY again each beaker and its contents. The loss of weight in the system is taken as the weight of carbon dioxide gener- ated. Make at least two determinations of this quantity. B. Second Step To determine the volume of the gas generated by the action of d,n acid on a weighed quantity of calcium car- bonate : 81/2 Weigh out 10.00 grams of the calcium carbonate, wrap in a paper as in No. 81/ l5 and transfer without loss to the gas generator. Now, carbon dioxide, being soluble in water, can not be collected over water without loss. The difficulty is avoided by using an intermediate bottle, holding about four times as much as the collecting bottle. It is provided with a two-holed stopper, through one hole of which passes a bent glass tube, terminating close to the bottom of the bottle ; through the second hole passes a tube terminating just below the stopper, and to the outer end of this is attached a rubber tube carrying the delivery tube. The delivery tube of the generator is now connected with the first tube of the large bottle, and the delivery tube of the latter is inserted in the mouth of the collecting bottle. The gas from the generator will thus be passed to the bot- tom of the large bottle and, being considerably heavier than air, will lie there some little time, while an equal volume of air will pass out from the top and may be collected arid measured, as in the hydrogen experiment. After the con- nections are made the apparatus should be tested for leak- age (see Exp. 41/ 3 b). Pour about 40 grams (30 c. c.) of nitric acid, a little at a time, through the thistle-tube until the reaction ceases. Note the volume of the acid used (why?), the temperature of the water in the tank, that of the room, and the height of the barometer. Make thus at least two determinations of the volume of gas generated, COMBINING WEIGHTS AND SPECIFIC PROPERTIES 51 Calculations. From these observations calculate first 81/3 the weight of one liter of the gas at and 760 millime- ters, then the specific gravity referred to hydrogen, and finally the ratio of the specific gravity to the combining weight. NOTE. The intermediate bottle, after once serving, 81/4 before being again used, must be cleaned of its gas. This may be done by blowing it out with a pair of bellows, the nozzle of which, extended by a rubber tube, reaches to the bottom of the bottle. EXPLANATORY NOTE. In this reaction the factors are calcium car- 81/5 bonate (a salt) and nitric acid ; the products are calcium nitrate (a soluble salt), water, and carbon dioxide, the latter appearing as a gas. This may be expressed by equation, thus : CaC0 3 4- 2HN0 3 = Ca(NO s ) 3 + H 2 -f C0 2 calc. nitric calc. water, carbon carb. acid. nitrate. dioxide. Recall the similar reaction in the experiment under Law 1, Chapter II (No. 36). 2. The Law of Dulong and Petit Relation between Equivalent and Combining Weights of Elementary Solids and their Specific Heats Definition of specific heat. The specific heat of a sub- 95 stance is the quantity of heat required to raise in tempera- ture 1 gram of the substance 1, divided by the quantity of heat required to raise 1 gram of water 1. It is necessary to discriminate carefully between quan- tity of heat and temperature. The unit of quantity (see No. 50/i) is called the gram-calorie, and is the quantity of heat required to raise one gram of water one degree (strictly from to 1, practically from 19 to 20 or thereabouts). Therefore specific heat may better be defined as the quan- tity of heat, measured in gram-calories, required to raise one gram of the substance one. degree. 52 ELEMENTARY PRINCIPLES OF CHEMISTRY 95/1 Specific problem. To determine experimentally the spe- cific heat, and to investigate the numerical relation be- tween this value and the combining weight in the case (A) of lead, (B) of zinc, (C) of tin.* 95/2 Specific heat of lead. Weigh out carefully 100.0 grams of lead (shot), transfer it to a large dry test-tube, set the tube and contents in water contained in a suitable vessel, and boil the water. The lead should be immersed below the surface of the water. Weigh out carefully in a beaker 50.0 grams of water which has the temperature of the room; test with the thermometer to see that the temperature does not change in an interval of five minutes. Put the thermometer, first drying it, in the lead, and see that it reaches a maximum which it holds for at least five minutes. Eecord this tem- perature. Then put the thermometer back in the water, stir, read the temperature carefully, and record it. Now, as quickly as possible, take the test-tube from the water, pour the warm lead into the water which the beaker contains, stir gently with the thermometer, quickly read the maximum temperature and record it. The observations therefore are : The weight of the water and its temperature just before mixing ; the weight of the lead and its temperature just before mixing ; the tempera- ture of the resultant mixture. 95/3 Calculations. Let x = the quantity of heat required to raise 1 gram of lead 1 ; y = the quantity of heat required to raise 1 gram of water 1 ; t = the degrees of temperature lost by the lead ; t'= the degrees of temperature gained by the water. * One of these three determinations may be thought sufficient, or perhaps all may be omitted if the student has already made similar ones in physics. COMBINING WEIGHTS AND SPECIFIC PROPERTIES 53 It follows that 100 X t X x = the quantity of heat lost by 100 grams of lead in cooling t. 50 X t' X y = the quantity of heat gained by 50 grams of water in warming t' . If now the experimental conditions are made such that the lead shall lose no heat save what goes to the water, and the water gain no heat save what comes from the lead, then these two quantities of heat must be equal. Equating the two expressions gives 100 X t X x = 50 X t' X y. From which is obtained - = Specific heat by definition. The experimental conditions called for to justify this equation can be but crudely realized by the directions here given, but sufficiently to give a good illustration of method. Calculate the quantity of heat required to raise 205 grams of lead 1. B Specific heat of zinc. Weigh out carefully 50.0 grams of 95/4 granulated zinc, and 50.0 grams of water. The procedure and calculations are the same as in the case of lead. Cal- culate the quantity of heat required to raise 65 grams of zinc 1. C Specific heat of tin. Use 50.0 grams of tin in coarse 95/5 powder, and 50.0 grams of water, and proceed as in B. Calculate the quantity of heat required to raise 118 grams of tin 1. NOTE. Save the material used in these experiments, so that it may be dried and be ready for use again. 21 54 ELEMENTARY PRINCIPLES OF CHEMISTRY 3. The Law of Mitscherlich Relation betiveen Composition, and hence Combining Weight, and Specific, i. e., Crystalline Form 107 No experiments. 4. The Law of Baoult (I) Relation between Combining Weights of Solutes and Specific Depressions of the Freezing Point in Specified Solvent (See Part I, Nos. 21 and 23/0 111 Specific problem. I. To investigate experimentally the relation between the depression of the freezing point and the quantity of the solute (A) when the solute is camphor and the solvent paraffin ; and (B) when the solute is naphthalene and the solvent paraffin. II. To investigate the numerical relation between the specific depression (i. e., depression for 1 gram of solute in 100 grams of solvent) of camphor in paraffin, and of naphthalene in paraffin, and the combining weights of the solutes. A 111/1 (1) Weigh out carefully 10.00 grams of paraffin and 0.80 of a gram of camphor. Melt the paraffin in an evaporating dish on the water-bath, then add the camphor and let it dissolve, stirring well. Pour a suitable portion of the solu- tion into a clean, dry test-tube. Let the liquid be about one inch deep. In another tube put some paraffin. Have a stirring-rod in each tube, and attach the tubes by rubber bands to a thermometer. Suspend the tubes in a beaker containing water, so that they are suitably immersed, and apply heat until the mixture melts. Allow to cool very slowly, keeping the mixtures and the water well stirred, and observe the freezing points. Repeat several times for con- stant results. It will save time to prepare the solution COMBINING WEIGHTS AND SPECIFIC PROPERTIES 55 called for in Exp. lll/ 2 , and to attach the three tubes at one time to the thermometer. (2) Make a mixture of 10.00 grams of paraffin and 1.20 111/2 grams of camphor, as in (1), and in a similar manner deter- mine its freezing point. Are the depressions approxi- mately proportional to the quantities of the solute in a constant quantity of the solvent ? Calculate the specific depressions i. e., depressions for 1 gram of solute in 100 grams of solvent and take their average. B (3) Make a mixture of 10.00 grams of paraffin and 0.60 111/3 of a gram of naphthalene, as in A ; (4) also of 10.00 grams 111/4 of paraffin and 1.20 grams of naphthalene. Observe the freezing points, and make the calculations as before. II. Multiply the specific depressions of camphor in 111/5 paraffin by the combining weight of camphor (C 10 H 16 = 151), and the specific depression of naphthalene by the combining weight of naphthalene (C 10 H 8 = 127), and com- pare the products. Are they approximately constant ? NOTE. It will be realized that the probable error of reading on the 111/6 ordinary thermometer is so large a fraction of the depression itself, that the results can not be quantitatively very satisfactory. A ther- mometer reading at least to hundredths would be needed for good results. Calculate the variation in the product made by the variation of one quarter of a degree in the depression. NOTE. The tubes may be cleaned by immersing in boiling water, 111/7 and the dishes by wiping with filter paper while still warm. 5. The Law of Baoult (II) Relation between Combining Weights of Solutes and Specific Elevations of Boiling Temperature in Specified Solvent (See Part 1,'^os. 21 and 24/ 4 ) Specific problem. I. To investigate experimentally the 127 relation between the elevation of the boiling temperature 56 ELEMENTARY PRINCIPLES OF CHEMISTRY and the quantity of the solute (A) when the solute is sodium acetate and the solvent water, and (B) when the solute is potassium tartrate and the solvent water. II. To investigate the numerical relation between the specific elevation of boiling temperature i. e., elevation for 1 gram of solute in 100 grams of solvent by sodium acetate in water, and by potassium tartrate in water, and the combining weights of the solutes. A 127/1 (1) Use a small distilling flask, as in the observation of boiling point (see Appendix, 17). Put in it a suitable quantity of water and some bits of pumice (why the lat- ter ? See No. 24/ 2 , Part I), boil, and observe carefully the temperature of the boiling water. Repeat all readings for constant results. 127/2 (2) Drain the water thoroughly from the flask, and put in its place a solution of 100.0 grams (or 100 cubi$ centi- meters) of water and 15.00 grams of sodium acetate with clean pumice. Again observe the temperature of the boil- ing solution. 127/3 (3) In a similar manner observe the boiling tempera- ture of a solution containing 100.0 grams (or 100 cubic centimeters) of water and 30.00 grams of sodium acetate. Are the elevations approximately proportional to the quantities of the solute in 'a constant quantity of the sol- vent ? Calculate the specific elevations, and take their average. B 127/4 (4) Observe the boiling temperature of a solution of 100.0 grams (or 100 cubic centimeters) of water and 20.00 127/5 grams of potassium tartrate; (5) also of 100.0 grams (or 100 cubic centimeters) of water and 40.00 grams of potas- sium tartrate. Make the same calculations as before. COMBINING WEIGHTS AND SPECIFIC PROPERTIES 57 II. Multiply the specific elevation of sodium acetate 127/6 (NaC 2 H 3 2 ) in water by the combining weight of the same, and the specific elevation of potassium tartrate (K 2 C 4 H 4 6 ) by its combining weight, and compare the products. Are they approximately constant ? The same remark as to accuracy of observation applies here as under the preceding law. Care should be taken that the bulb of the thermometer is completely immersed in the liquid. NOTE. The solutions may be saved in order to recover the salts by crystallization. Alternative experiment Instead of the solutions indi- 127/7 cated in the preceding experiment, may be used the follow- ing : A 15.40 per cent and a 30.80 per cent solution of potas- sium chloride (KC1), and an 11.00 per cent and a 22.00 per cent solution of ammonium chloride (NH 4 C1). CHAPTER VI EXPERIMENTS ILLUSTRATING THE METHOD OP DETER- MINING EQUIVALENT AND COMBINING WEIGHTS OF ELEMENTS AND FORMULAS OF COMPOUNDS 1. Determination of the Equivalent Weight of an Element 144 Specific problem. To determine experimentally the mass of tin which combines with 7.94 grams of oxygen. (Tin does not combine with hydrogen.) 14-4/1 EXPLANATORY NOTE. It is not practicable to cause a weighed quan- tity of tin to combine with oxygen of the air, and to weigh the oxide obtained ; but by the action of nitric acid, the oxidizing power of which has already been noted [in what connection f], an oxide of tin is produced which does not volatilize, does not combine with nitric acid, and does combine with water, forming a compound, however, from which the water is expelled by prolonged heating, leaving simply the oxide. The excess of nitric acid is also volatilized by heat. 144 Weigh very carefully to hundredths of a gram a small evap- orating dish, clean and dry, together with a short glass rod. Weigh out carefully 5.00 grams of pure tin foil. Put a portion of this in small pieces in the evaporating dish, and just moisten with nitric acid. When the action is nearly over, put in more tin, and then moisten again with acid. Continue this until all the tin is used, with the minimum of acid that will suffice. Describe what takes place. What is the brown gas produced? Have care not to inhale it ; it is well to use the hood. Also bear in mind the extreme corrosiveness of nitric acid. Use the rod to pour by, leaving it in the dish. Keep the outside of the beaker clean. Have a wet sponge at hand. 58 DETERMINING EQUIVALENT WEIGHTS 59 Heat the dish and contents on the water-bath until the liquid has entirely volatilized, transfer to the iron ring of the stand, apply the direct flame, gently at first passing it back and forth, then give it the full heat of the burner, continuing for some time after the fumes have ceased to appear. The powder should be slightly brown when hot, and yellowish when cool. Let the dish become thoroughly cool, and weigh it with its contents to hundredths of a gram. To insure the com- pleteness of the operation, ignite again for fifteen or twenty minutes, let cool, and weigh again. Assuming that you have in the dish a compound con- 144/2 taining only tin and oxygen, calculate the weight of oxy- gen which has combined with the 5 grams of tin, and from this the weight of tin which combines with 7.94 grams of oxygen. Make at least two complete determinations. 2. Determination of the Combining Weight of an Element Under the law of Dulong and Petit choose what mul- 150 tiple of the equivalent weight determined in No. 144/ 2 shall be taken as the combining weight of tin. The specific heat of tin is 0.056. (See Nos. 101 and 105, Part I.) It will be realized that the utmost care must be taken 150/1 in the manipulation, since the weighings must be to hun- dredths of a gram, and the variation in the quantity of oxygen for 5 grams of tin should not be more than one or two hundredths of a gram. Having determined the combining weight of tin, deduce the for- 150/2 mula for the oxide. Try to write the equation for the reaction between tin (symbol, Sn) and nitric acid, HNO S . What is the percentage of tin and of oxygen in tin oxide ? 3. Determination of the Formula of a Compound Specific problem. To determine the percentage of the 154 proximate constituents, carbon dioxide, C0 2 , and sodium 60 ELEMENTARY PRINCIPLES OF CHEMISTRY oxide, Ka.,0, in the compound, sodium carbonate, and from these to deduce the formula of the compound. 154/a (a) To determine the percentage of carbon dioxide : Take a sufficient quantity of the sodium carbonate for all the following experiments, heat in the larger evaporating dish over the flame, gently and not too hot. This is simply to secure the dryness of the sample. Weigh out carefully 5.00 grams of the dry sample, and follow the procedure given for determining the weight of carbon dioxide in calcium carbonate (see Exp. 81/j), using hydrochloric instead of nitric acid. Make at least two determinations. Calculate the result as grams of carbon dioxide in 100 grams of the sample. Secondary observation. Filter and crystallize the salt which is left in solution. What substance is it ? (I) To determine the percentage of sodium oxide : 154/1 EXPLANATORY NOTE. Since it is not practicable to separate this constituent from the compound, and weigh it by itself, it is necessary to convert it quantitatively into some substance which it is practicable to separate and weigh, and the composition of which is known. For this purpose the carbonate is converted into sodium nitrate by the action of nitric acid, the carbon dioxide escapes, and the water and excess of nitric acid are volatilized. Sodium nitrate is not volatile, and is not decomposed by heating short of fusion. Its composition is given by the formula Na 2 N 2 06, or, as usually expressed, NaNOs. [What per- centage of Na 2 does this formula show ?] 154/b Weigh out carefully 5.00 grams of the dry sample of sodium carbonate in a small evaporating dish, the weight of which has been determined to hundredths of a gram. Neutralize this with nitric acid, using a very slight excess, and having care not to lose by effervescence. Evaporate on the water-bath until the liquid has disappeared, and then, passing the burner back and forth, heat gently until the substance just begins to melt. Let cool, and weigh. The substance obtained is sodium nitrate. Make thus at DETERMINING EQUIVALENT WEIGHTS 61 least two determinations of the weight of this substance obtained from 5 grams of the carbonate. Knowing the percentage of sodium oxide in sodium nitrate, calculate the weight of sodium oxide contained in the weight of the nitrate obtained, and reckon this result as grams of the oxide in 100 grams of the carbonate. Secondary observation. Redissolve and crystallize the sodium ni- trate. (c) To deduce the formula : Since the proximate constituents are carbon dioxide, 155 C0 2 , and sodium oxide, Na 2 0, the formula must be (Na 2 0)a; (C0 2 ) y , the coefficients, x and /, to be deduced from the percentages experimentally obtained. Hence divide the percentage of each constituent by the combining weight of the same, and the quotients give the ratio of the coefficients, x and y. Now these quotients are not necessarily whole numbers, but the peculiarity of chemical phenomena is that they stand in the ratio of whole numbers, usually quite small. Hence inspection or division by the smallest will suffice to show the smallest whole numbers that have the ratio of the quotients. In this specific problem the ap- proach of the quotients to equality measures the accuracy of the experimental work. The simplest values, therefore, for x and y are 1 and 1, and, in the properties of this sub- stance, no reason is found for using any multiples of these values. The accepted formula is therefore Na 2 OC0 2 , or, as it may also be written, ]STa 2 C0 3 . [Write the equation for the reaction between sodium carbonate and nitric acid.] CHAPTER VII THE ATOMIC THEORY 157 to N O experiments. 187 CHAPTER VIII RELATION BETWEEN THE PROPERTIES OF THE ELEMENTS IN GENERAL AND THEIR COMBINING WEIGHTS Experimental Study of the Properties of the First Twenty- five Elements (in the Order of their Combining Weights] and Some of their Compounds So far as practicable, each element in turn will be presented in its free condition for descriptive study : first as to its physical properties, those which appear on inspection, and others ; second, as to its chemical properties. Then will follow the study of some of its most important compounds. It is judged unnecessary to give detailed directions for manipulation in all experiments, as by this time the student should have had sufficient experience in the laboratory to give him some judg- ment of his own as to how to do things. Some of the manipulations, too, are repeated many times. 1. HYDROGEN Symbol H. Comb. wt. 1 202 Preparation. See Exps. 17/i and 41/ 4 and Appendix, 19, I. Use the small collecting bottles or test-tubes. Use iron (nails) or zinc and hydrochloric or dilute sulphuric acid. NOTE. Hydrogen with, air makes a mixture which may explode dangerously on ignition. Therefore the greatest 62 DESCRIPTION OP ELEMENTS AND COMPOUNDS 03 caution must always ~be taken to avoid accident from unex- pected ignition by contact with flame. Physical Properties Observe as to color, odor (any odor is due to impurity). 203 As to weight compared with air. A crude test : Fill a test- tube with hydrogen over water, close with the thumb, hold its mouth downward directly over and close to the mouth of a second tube, then remove the thumb quickly, bringing the two ends together, reverse the position of the tubes, separate them, closing the second, now the upper one, with the thumb, and test at the flame for the presence of hydro- gen in both tubes. As to solubility in water. A crude test: Collect in a 204 test-tube, leaving a little water; close tightly with the thumb, shake, open under water, close, and shake again. Eepeat this several times, and note if the volume of water in the tube increases and the gas decreases. As to relative divisibility. Use a dry glass tube, closed 203/1 at one end with a porous plug of plaster of Paris, and open at the other. Close temporarily the plugged end with the thumb or a cork, and fill the tube thoroughly by dry dis- placement i. e., hold the tube open end downward, thrust the delivery tube at first well up to the top, then slowly withdraw it (compare Appendix, 19, VI). When the tube is full, immerse the lower, open end in water, uncover the porous plug, and allow to stand for some time. What takes place ? Can you explain the phenomenon ? A modification of the preceding experiment, suitable to be performed 203/2 by the teacher before the class : Use a porous battery jar, closed by a cork or a plaster plug through which passes a glass tube of convenient length. Support this so that the lower open end of the tube is immersed in a col- ored liquid (e. g., water, colored by potassium permanganate). Slowly lower over the porous jar a glass bell jar filled with hydrogen. When the gas ceases to bubble from the end of the tube, remove the bell jar. The same may be tried with illuminating gas in place of hydrogen. 64 ELEMENTARY PRINCIPLES OF CHEMISTRY 205 Occlusion. Hold a piece of platinum sponge (i. e., finely divided platinum deposited on the surface of asbestos fiber), thoroughly dried, in the stream of hydrogen issuing from the delivery tube, but only when the gas is sufficiently pure to be ignited with safety. [This experiment, like the preceding one, may best be performed by the teacher.] Chemical Properties Test as to the action of hydrogen on wet litmus paper. 207 Test as to combustibility. With suitable precaution, the gas issuing from the delivery tube may be ignited as a jet, but, owing to the danger of explosion, this should not be tried until the action has continued long enough to expel the air from the generator. It is best always to test a small sample of the gas thus : Collect in a test-tube over water, close with the thumb, bring it to the flame, mouth down- ward, remove the thumb, ignite the gas, and slowly invert the tube. The gas should burn quietly in the tube. If it burns quickly with a slight report, it is not safe to ignite at the generator. 207/1 Make it a rule always to wrap several folds of the towel over and around the generator before igniting the gas, no matter hoiv confident you may be of its purity. Serious ac- cident may follow neglect of this precaution. Observe the character of the hydrogen flame ; hold a dry glass plate or beaker over the flame. What is the pro- duct of combustion ? 208 Nascent state. Use four test-tubes, fill about two thirds full with water, and color each portion slightly with a few drops of potassium permanganate solution. Into the first put a few fragments of zinc and a few drops of sulphuric acid, and let the action continue for some minutes. What is the result ? Whatsubstance is it which causes the color to disappear? To answer this,. place in the second tube some zinc, in the third some sulphuric acid, and into the fourth let the hydrogen from the generator slowly bubble, preferably using zinc as the metal. DESCRIPTION OF ELEMENTS AND COMPOUNDS 65 EXPLANATORY NOTE. The permanganate loses its color in conse- 208/1 quence of losing some of its oxygen. The oxygen is taken from it by the hydrogen, but only at the moment of its liberation from the acid, when it is said to be in the nascent state. The experiment may not be entirely satisfactory, for the reason that the impurity which is usually carried by the hydrogen from the generator, and which comes from the metal, and imparts the odor to the gas, tends to remove oxygen from the permanganate. Observe as to the odor of the gas which has bub- bled through the colored liquid. To explain the phenomenon of the nascent state it has been sug- 208/2 gested that the substance in the atomic condition may show an activity which is wholly or partly lost after the atoms have come together into molecules. Another method of generating hydrogen. Use a few frag- 202/3 ments of zinc in a test-tube with a dilute solution of sodium hydroxide. The reaction is facilitated by a piece of iron (a nail) in contact with the zinc. 2. LITHIUM Li.-6.97 No experiments, unless to show a sample of some 210 salt of lithium e. g., the chloride and the color it im- parts to the flame. 3. GLUCINUM or BERYLLIUM Gl. or Be. 9.0 No experiments. 216 4. BORON B. 10.86 Boron is difficult of preparation, so it is impracticable to show it. It forms an oxide whose symbol is B 2 3 , and which combines with water, forming boric acid. The sodium salt of this is the familiar com- mercial substance, borax. Borax and boric acid crystallized. Saturate about a 225 beakerful of water with borax. Filter, if not clear, and 66 ELEMENTARY PRINCIPLES OP CHEMISTRY set aside about one half of this to crystallize. To the re- mainder add about one half its volume of hydrochloric acid, and set this aside to crystallize. Observe the different ap- pearance of the two substances when crystallized. Kinse a sample of each substance quickly with a very little cold water, and dry it on filter paper. Observe the reactions on litmus paper. 226 Flame coloration. Put a very small quantity of borax in an evaporating dish, add a few drops of alcohol, ignite the alcohol, and observe the color of the flame, while stirring the mixture. Extinguish the flame, add a few drops of dilute sulphuric acid, reignite the alcohol, and again ob- serve the flame color. This test is often used to recog- nize boric acid and its compounds. 228 The borax bead. Melted borax has the property of dis- solving many metallic oxides, and these often impart charac- teristic colors to the substance. This test is very useful in qualitative analysis. To make a borax bead, use a piece of platinum wire, about 75 millimeters (3 inches) long, one end of which is fused into a piece of glass tubing for a handle. Bend the free end of the wire around the sharp- ened end of a lead pencil, so as to make a small loop. Heat this red hot, touch it to some powdered borax, and heat the borax which adheres to the wire. What is the first effect of this, and to what is it due? (see Exp. 21/ 7 .) When the borax has fused to a clear, glasslike bead, touch it, while still hot, to some particles of copper oxide. These adhere to the hot bead. Do not let it take up too much a very minute quantity suffices ; and if you have not enough on first trial, it is easy to add more, but it is not so easy to reduce the quantity. Heat again to fuse the bead with the adhering copper oxide. The particles of the latter can be seen slowly to dissolve, often with beautiful play of colors on the surface of the bead. Observe the color of the bead, when hot, and when cold. The bead should be transparent, but, if too much material has been put in, it DESCRIPTION OF ELEMENTS AND COMPOUNDS 67 will be opaque, so the color can not be distinguished. In this event, clean the wire, leaving only a fragment of the bead, and add more borax. To clean the wire, heat the 229 bead hot, and quickly cool it in water. This makes the mixture brittle, so that it can be easily removed. To get the best results, in the bead experiments, the mouth blow- pipe should be used for heating. By this means the char- acter of the flame can be so varied as to cause the oxida- tion of the dissolved substance, or the removal of oxygen from it i. e., reduction. These changes often produce changes in color. Thus the copper bead is blue when heated in the oxidizing flame, and it has the color of metal- lic copper in the reducing flame. (Concerning the use of the blowpipe see Appendix 21.) Manganese in the bead. After seeing the color of the 228/2 copper bead in the oxidizing and in the reducing flame, make a new bead and color it with manganese dioxide. Ob- serve the color in each flame. 5. CARBON C. 11.91 Elementary carbon exists in three allotropic forms, 234 diamond, graphite, and charcoal ; only the last is here studied. Preparation. Heat a small quantity of sugar on the 248 spatula blade. It burns and chars. The black charred material consists mainly of carbon. In similar manner many other substances show the charring effect of heat, and this is evidence of the presence of carbon as a con- stituent in the substance. It is due to the removal in part of other constituents, particularly hydrogen, by combustion, the less easily combustible carbon being left as residue. Place some fine shavings in a dry test-tube, fitted with 254 a cork and a short glass delivery tube, and apply the full heat of the flame. Note the appearance of moisture and 68 ELEMENTARY PRINCIPLES OF CHEMISTRY white fumes. Ignite the gas issuing from the tube. After the heating observe the black residue in the tube. This is charcoal. Note also the liquid, having tarry odor. Test it with litmus. The application of heat in this manner, without free access of oxygen or air, is called dry or destructive distil- lation. It is applied on a large scale to the production of various substances from wood, and likewise to the produc- tion of illuminating gas and other substances by the de- 253 structive distillation of coal. Heat a little soft coal in a test-tube, as in No. 254. Properties of Carbon ( Charcoal) 254/1 Porosity. Holding a piece of charcoal in the tongs, plunge it beneath the surface of some hot water. Is char- coal heavier or lighter than water ? Keep it under hot water for a considerable time. Does it finally sink in the water ? 255 Absorptive power. Generate a very little hydrogen sul- phide, using a piece of iron sulphide not larger than a bean and a few drops of hydrochloric acid, and collecting the gas in a dry bottle covered with a glass plate (Appendix, 19, V). Drop into this a piece of charcoal, previously heated for a few seconds in the gas flame. Set this aside with the cover on, and some time later observe by odor if the gas has been absorbed by the charcoal. 255/1 In a beaker containing water, colored by a few drops of potassium permanganate solution, place a small quantity of boneblack (animal charcoal), boil for a few minutes, then filter. Does the filtrate lose its color ? Is the color re- moved by filtering simply, without the use of the charcoal ? 257 Combustion. Does the substance burn? Does it burn with a flame ? Why not ? 257/1 Product of combustion, Ignite a piece of charcoal in the flame until it is well aglow, then drop it into a bottle and cover the latter with a glass plate. When the coal is ex- DESCRIPTION OF ELEMENTS AND COMPOUNDS 69 tinguished, either take out the remaining charcoal or invert the bottle over another bottle, pouring thus the heavy gas from the first into the second bottle and leaving the char- coal. Next pour into the second bottle some lime-water (calcium hydroxide solution, Ca0 2 H 2 ) and shake the con- tents. The white precipitate is calcium carbonate, CaC0 3 , and its formation may generally be taken as evidence of carbon dioxide, which is, in fact, the combustion product of carbon. Write the equation for the formation thus of cal- cium carbonate. Reducing power. Make a shallow hole in a piece of 257/2 charcoal, place in this a small quantity of litharge (lead oxide), and turn the flame (reducing) of the blowpipe upon this (see Appendix, 21, II). Describe and explain what takes place. 5a. Carbon Dioxide, C0 2 Preparation. Use calcium carbonate and hydrochloric 261 acid in the generator. Collect the gas by upward displace- ment of air (Appendix, 19, Y) in the small collecting bot- tles. As to solubility. Test as in Exp. 204. 262 As to combustion and specific gravity. Plunge the lighted 263 taper slowly into a bottle filled with the gas. Place the lighted taper in one bottle, and pour upon it the heavy gas from a second bottle. Holding a piece of magnesium ribbon in the tongs, ignite it at the flame and plunge it into a bottle well filled with the gas. What is the white powder seen in the bottle after this experiment ? What is the black substance ? Why does it retain some- what the form of the ribbon ? As to acid reaction. Test the action on wet blue litmus 264 paper. As to reaction with lime-water (calcium hydroxide solu- 265 tion). Collect some of the gas in a test-tube, pour in some lime-water, and shake. Pass in repeatedly more of the gas 22 70 ELEMENTARY PRINCIPLES OF CHEMISTRY until the precipitate at first formed nearly or quite disap- pears. Then boil the contents of the tube. Explain the reprecipitation. 265/1 EXPLANATORY NOTE. Excess of carbon dioxide in water forms with calcium carbonate an acid carbonate which is soluble in water. This compound is broken up at the temperature of boiling and the ex- cess of carbon dioxide is expelled, in consequence of which the insol- uble calcium carbonate reappears. Many natural waters contain this acid carbonate of calcium or of other bases, and therefore become tur- bid and deposit sediment by boiling. Test some sample of natural water for this effect. See also Part I, No. 271. * 267 Carbon dioxide in the breath. Through a glass tube blow air from the lungs into some li le-water. 5d. The Hydrocarbons Methane, CH 276 Preparation. Take about equal bulks of lime, CaO, and of sodium acetate, ^N"aC 2 H 3 2 ; pulverize and mix them thoroughly in the mortar ; put some of the mixture in a dry test-tube (about three quarters full), which is provided with a cork and delivery tube ; support this on the iron ring, heat, and collect the gas over water in test-tubes. 277 Properties. Examine the gas as to color, odor, solubil- ity in water, and action on litmus (after the sample of gas has been shaken with w NaOH + A1 2 (S0 4 ) 3 = A1 2 6 H 6 NaOH + NH 4 N0 3 = NH 4 OH + NaOH + NaN0 3 = no reaction. 10. MAGNESIUM Mg. 24.10 445 The element itself is supplied. Describe its appear- ance and recall whatever has been learned of it in previous experiments. By observation answer the following ques- tions : Does magnesium melt ? Does it burn ? What is 446 the appearance of the product of combustion, MgO ? Does the oxide dissolve in water ? In hydrochloric acid ? Using strips of the ribbon, one or two centimeters long (one half to three quarters of an inch), ascertain if magne- sium decomposes cold water or boiling water. (Have a DESCRIPTION OF ELEMENTS AND COMPOUNDS 81 piece of red litmus paper in the water.) Does the hydrox- ide act as base or acid ? Does the metal dissolve in very dilute hydrochloric acid ? In very dilute nitric acid ? In very dilute sulphuric acid ? Make a dilute solution of magnesium sulphate (MgS0 4 ), 446/1 by dissolving the dry salt (from the side table), a lump about the size of a pea, in a test-tube nearly filled with water. Add to a small portion of this solution a few drops of sodium hydroxide solution. What is the precipitate? (See No. 420/ 2 .) In a similar manner add to diiferent por- tions of the magnesium sulphate solution a few drops of solutions of these several substances : ammonium hydrox- ide, ammonium chloride, ammonium carbonate, sodium carbonate, sodium phosphate. Magnesium hydroxide, car- bonate, and phosphate are insoluble. With this informa- tion complete the following equations : Mg + = Mg +1H 2 = Mg -fjHCl .= Mg01 2 - Mg + H 2 S0 4 = MgS0 4 + * * S0 MgS0 4 -+fNaOH ?*&&* MgS0 4 + Na 2 C0 3 = MgC0 8 + MgS0 4 + *Ta 2 HP0 4 = MgHP0 4 What is the valence of magnesium ? 11. ALUMINIUM Al. 26.9 The element itself is supplied. Describe its appearance. 451 to By trial learn if it melts, and if it burns. Does it decom- 454 pose water, cold or boiling ? Does it dissolve in dilute hy- drochloric acid ? In dilute nitric ? In dilute sulphuric ? In dilute solution of sodium hydroxide (boil) ? What is the behavior of aluminium sulphate (alum) in dilute solu- tion toward these several substances also in dilute solution : 82 ELEMENTARY PRINCIPLES OF CHEMISTRY sodium hydroxide (first in small proportion then in excess), ammonium hydroxide, ammonium chloride, sodium carbon- ate, sodium phosphate ? Aluminium hydroxide and phos- phate are insoluble in water ; the chloride is soluble. The hydroxide dissolves in excess of sodium hydroxide, forming a soluble compound in which the former acts as acid, and the latter as base. 12. SILICON Si. 28.2 EXPLANATORY NOTE. It is not practicable to supply the element not, indeed, even to show it. The dioxide, Si0 2 , is, however, abundant 473 as common sand. This oxide acts as an acid-former, and when brought in contact with fused sodium carbonate it liberates carbon dioxide and combines with the bafte, forming a silicate, the composition of which may be sufficiently indicated by the formula Na 2 OSi0 2 or Na 2 Si0 3 . This is soluble in water, and when hydrochloric acid is added to the solu- tion the silicic acid, H 2 OSi0 2 , in turn is liberated and, being insoluble in water, is precipitated. The precipitate is, however, soluble in the hydrochloric acid, probably without chemical change, and may be repre- cipitated by neutralizing the acid, HC1, with ammonium hydroxide. 474 Place a few grains of sand (powdered glass may be used in the same way) in a shallow hole of a piece of charcoal, cover the sand with five or six times its bulk of sodium carbon- ate, fuse thoroughly, using the blowpipe. Note the minute bubbles. What causes them ? After sufficient fusion, re- move the mixture from the charcoal, treat with hot water, filter, acidulate with hydrochloric acid, using litmus paper, and boil. If the precipitate does not appear, add ammo- nium hydroxide to slight alkalinity, and boil again. De- scribe the appearance of the precipitated silicic acid, H 2 Si0 3 . Filter the precipitate and dry it. Complete the equations : Si0 2 + Na 2 C0 3 = NapSiOs + Na 8 SiOs -4/HQ1 = H 8 SiO s + * ( ' H 3 Si0 3 (heated) = SiO g + ^ c DESCRIPTION. OF ELEMENTS AND COMPOUNDS 83 13. PHOSPHORUS P.-30.8 Recall the use already made of phosphorus. It must be handled with the utmost caution, because of the danger of ignition. Burns produced by it are painful and difficult to heal. Describe the appearance of yellow phosphorus. 486 Place the porcelain lid and cork on a dry glass plate. 487 On the lid place a small fragment of phosphorus. Ignite the latter and cover it with a dry beaker. The oxide formed is mainly Pg0 5 . Describe the burning and the product. Let the white oxide settle upon the glass plate. Observe the marked deliquescence. Touch blue litmus paper to the oxide. Is it an acid- or a base-forming oxide ? Phosphorus a reducing agent, Into a test-tube about one 487/1 third filled with a dilute solution of silver nitrate, AgN0 3 , drop a small piece of yellow phosphorus, about twice the size of a pinhead. Le1^ stand for twenty-four hours. Minute crystals of metallic silver will be seen at the bottom of the test-tube. The liquid may be drained off and the crystals fused on the charcoal or asbestos to a small globule of silver. Describe the appearance of the red phosphorus, and 490 compare it with the yellow as to inflammability. EXPLANATORY NOTE. Phosphorus pentoxide combines with water 494 to in three different ratios, forming three phosphoric acids, thus : 496 P 2 B + 3H 2 = 2H 3 P0 4 orthophosphoric acid. P a 6 + 2H 2 = H 4 P 2 07 = pyrophosphoric acid. P 2 6 + H 2 = 2HP0 3 = metaphosphoric acid. When sodium orthophosphate, Na 2 HP0 4 , is sufficiently heated, it be- comes the pyrophosphate, Na 4 P 2 7 . Both are soluble in water, but the latter does not immediately pass back to the orthophosphate when in contact with water. Taking a fragment of crystallized sodium orthophos- phate, Na 2 HP0 4 , fuse it on charcoal, keeping it fused for some minutes. Eemove the salt when cool, dissolve in a 84 ELEMENTARY PRINCIPLES OF CHEMISTRY little hot water, filter, and to the solution add a few drops of silver nitrate solution. Describe the precipitate, Ag 4 P 2 7 . Dissolve a small piece of the sodium orthophosphate in water, and add to this a few drops of silver nitrate. De- scribe this precipitate, Ag 3 P0 4 . How has the fusion changed the orthophosphate ? Write the equation for the reaction of fusion, also for the precipitations. 14. SULPHUR S.-31.83 509 Sulphur was studied in considerable detail at the very first of the course, and it would seem unnecessary to repeat the experiments here. They should at least be recalled, and it may be well to write out a de- scription from the earlier observations. These items may serve as out- line : specific gravity, as to solubility in water, color, hardness, crys- talline form from fusion, from solution in carbon disulphide, plastic allotropic form, fusion, combustion, behavior with iron, with zinc, and with lime. 510 Hydrogen sulphide, H a S, also was studied somewhat, and these items and should be recalled : formation by the action of hydrochloric acid on 511 metallic sulphides, odor of gas, action on paper wet with lead acetate, behavior with strong nitric acid. 512 The sulphides. Hydrogen sulphide is an acid, and in combining with ammonium hydroxide it forms the acid sulphide, NH 4 HS, or the normal (NH 4 ) 2 S. Persulphides also are formed, such as (NH 4 ) 2 S 2 and (NH 4 ) 2 S 3 , the coeffi- cient of S being 2 or more. The common reagent, yellow ammonium sulphide, much used in the laboratory, is made by saturating ammonium hydroxide with the gas, H 2 S. It therefore contains more or less persulphide. With this solution of ammonium sulphide make the following experi- ments and describe the results : Add a small portion to these several solutions : hydrochloric acid, sodium nitrate, magnesium sulphate, alum, iron sulphate, copper sulphate, lead acetate, and, if at hand, to solutions containing arsenic, antimony, and cadmium. jtfr^ A p DESCRIPTION OF ELEMENTS ANJT COMg0T?BS)S r Complete the following equations: S + MgS0 4 = S + AL(SOJ 3 = A1 2 6 H 6 S + FeS0 4 = FeS S + CuS0 4 = CuS What is the valence of sulphur in hydrogen sulphide ? Sulphur dioxide. This substance is formed by burning 515 sulphur, but more conveniently by the action of hydro- chloric acid on sodium sulphite, Na 2 S0 3 . Place some of the latter in the generator, add the acid gradually, and collect the gas by dry displacement, since it is soluble in water and heavier than air. Test the gas for odor (it is 516 very irritating), solubility, action on the burning candle or match, weight as compared with air (as in the experiment with carbon dioxide Xo. 263), and action on litmus paper, both brief and somewhat prolonged. To a portion of sodium sulphite solution add a few 517 drops of potassium permanganate, KMn0 4 , and to another portion add a crystal of iodine. Describe what takes place. Sulphur dioxide, sulphurous acid, H 2 S0 3 , and the sulphites, all tend more or less readily to take up oxygen and become the trioxide, sul- phuric acid, H 2 S0 4 , and the sulphates respectively; the former there- fore act as reducing agents. Complete these equations : 517/1 H 2 S0 3 3 - = Na 2 S0 3 + = nw * ^ JSTa 2 S0 3 + H 2 4JI = 2HI +,)4 > @ 2 0/SWOv + H^T = 2KOH + 2MnO 23 ^ 86 ELEMENTARY PRINCIPLES OF CHEMISTRY 519 Sulphuric acid, With sulphur trioxide it is not practi- cable to work, but its compound with water, sulphuric acid, H 2 O.S0 3 , or H 2 S0 4 , is a common reagent. Describe its appearance. Taking about one half of a test-tubeful of the concen- trated sulphuric acid, dilute it with about three times its volume of water. Note what takes place. CAUTION : Make it a rule in diluting strong sulphuric acid always to add the acid to the water slowly and with constant stirring, and never the water to the acid. Why ? 521 Taking three equal portions of the dilute acid, neutral- ize one portion with a solution of sodium hydroxide, filter, concentrate, and crystallize. Neutralize the second por- tion in the same way, and to it add the third portion ; filter, concentrate, and crystallize. Compare the crystallizations ; dry a sample of each product on filter paper, and test it with litmus paper, red and blue. Let a sample of each lie exposed to the air for twenty-four hours. 519/1 To show the tendency of strong sulphuric acid to re- move the constituents of water : having a small quantity of the acid in a test-tube, drop into it a splinter of pine wood, or a piece of filter paper. What takes place after some minutes? 521/1 Sulphate changed to sulphide : taking a crystal of sodium sulphate from the side table or from the product of the preceding experiment, No. 521, fuse it on the char- coal, first adding some sodium carbonate (use the reducing flame of the blowpipe). Kemove the product from the charcoal, lay it upon a piece of filter paper wet with lead acetate, and let a drop of water fall upon it. What does the black stain indicate ? A silver coin may be substituted for the lead paper. 521/2 Solubility of sulphates. Add a few drops of sodium sul- phate or of dilute sulphuric acid to a solution of lead ace- tate ; also, if at hand, to the solution of a barium salt (e. g., BaCl 2 ), and of a strontium salMe. g. . '/ ' DESCRIPTION OF ELEMENTS AND COMPOUNDS 87 Complete these equations : S0 3 -fH 2 = H 2 S0 4 + NaOH = Ka 2 S0 4 + H 2 S0 4 + NaOH = H 2 S0 4 + BaCl 8 = BaS0 4 Na 2 S0 4 + Pb( A) 2 = PbS0 4 + 15. CHLORINE Cl.-35.18 This element is extremely disagreeable to deal with in large quantity. It is best that the illustrative experiments in the laboratory be only on the test-tube scale. Preparation. Hydrochloric acid is commonly the source 529 of chlorine. From this it is obtained by oxidizing the hydrogen to water, and thus liberating the chlorine. The methods differ only in the means used to oxidize. (a) Place in a test-tube a small quantity of manganese 529/2 dioxide, Mn0 2 , about as much as could be heaped on a cop- per cent. Add to this about one quarter of a test-tubeful of concentrated hydrochloric acid and warm gently. Ob- serve cautiously the odor, also the effect on wet litmus 530 paper of brief and then of somewhat prolonged exposure to the gas. Enough of the gas may be generated in the test-tube to show its color, and some of it may be poured into a second dry tube (it is heavier than air) in order to test as to its solubility in water. Complete this equation by inserting the necessary co- efficients : Mn0 2 + HC1 = MnCl 2 (sol. salt) + H 2 + 01. (b) Place in a test-tube some potassium chlorate, KC10 3 , 529/1 about the size of a pinhead, and add concentrated hydro- chloric acid, a few drops. Note the evolution of gas, the 88 ELEMENTARY PRINCIPLES OF CHEMISTRY odor, color, etc. Dilute with water. Does the reaction continue ? Complete the equation : KC10 3 + HC1 = KC1 (sol. salt) + H 2 + 01. Hydrochloric Acid, HCl 534 Preparation. Place some sodium chloride, NaCl, in a test-tube ; a lump about the size of a pea will serve. Add to this a few drops of concentrated sulphuric acid, H 2 S0 4 , and warm. Note the odor of the gas, and its action on blue litmus paper. Hold at the mouth of the test-tube a glass rod wet with ammonium hydroxide. What causes the white fumes ? The properties of hydrochloric acid have been illus. trated in its repeated use. 534/1 Solubility of chlorides. Mix a few drops of hydrochloric acid, or of sodium chloride solution, with these several salts in dilute solution : lead acetate, Pb(C 2 H 3 2 )2 ; silver ni- trate, AgN0 3 ; mercurous nitrate (if at hand), HgN0 3 ; magnesium sulphate, MgS0 4 ; iron sulphate, FeS0 4 , etc. The first three of these bases form insoluble chlorides. The other chlorides are as a rule soluble. Complete these equations : NaCl + H 2 S0 4 = HCl + Pb(C 2 H 3 2 ) 2 = HCl + AgN0 3 = HCl + HgN0 3 = HCl + MgS0 4 = 16. POTASSIUM 552 If the metal potassium and its hydroxide, KOH, are at and hand, experiments with them may be made exactly as with 553 sodium and sodium hydroxide. But the results are so en- tirely similar that the observations may be omitted. DISCRIPTION OF ELEMENTS AND COMPOUNDS 89 17. CALCIUM Ca. 39.7 The element is not available for these experiments, 581 being difficult to obtain, but the oxide, lime, CaO, is read- ily procured. If quicklime is at hand, use it to observe the effect of contact with water. If only slaked lime, Ca(OH) 2 , is provided, it will serve, except for this single observa- tion. Learn by experiment if calcium hydroxide dissolves 582 in water. Does it act as base or as acid ? What reaction takes place with carbon dioxide ? Prepare a solution of calcium chloride, CaCl 2 , by neutralizing with lime about one third of a test-tubeful of the acid, and filtering. With 584 this solution, observe the reactions with dilute solutions of these several substances : ammonium hydroxide, chloride, carbonate, and oxalate (NH 4 ) 2 C 2 4 ; sodium phosphate, and sodium or potassium sulphate or dilute sulphuric acid. Complete the following equations : . Qa(OH) 2 + SCI = Co CaCl 2 +(NH 4 ) 2 C0 3 - ^^3 ' CaCl 2 + (NH 4 ) 2 C 2 4 = j <* IU * CaCl 2 -fNa 2 HP0 4 = CaCl 2 -f]Sra 2 S0 4 = 23. IRON Fe. 55.6 Several experiments with iron and its salts have already 596 been performed, and should be recalled here. With a sam- ple of iron, learn by trial if it melts or burns in the flame of the Bunsen burner or in that of the blowpipe. Does it decompose water, cold or boiling ? Does it dissolve in 598 dilute hydrochloric acid (recall other experiments) ? In dilute nitric acid ? In dilute sulphuric acid ? 90 ELEMENTARY PRINCIPLES OF CHEMISTRY Take about a test-tubeful of hydrochloric acid, dissolve in this as much iron as it will take up, keeping some ex- cess of iron present (a nail serves well). Filter, concen- 600 trate, and crystallize the ferrous chloride. As alternative with the preceding, dilute sulphuric acid may be used and the ferrous sulphate may be crystallized. To a dilute solution of ferrous salt freshly made, add ammonium hydroxide solution. What is the precipitate ? 600/1 To a dilute solution of ferrous salt add a few drops of concentrated nitric acid, boil a few seconds, then add am- monium hydroxide to alkaline reaction. What is the pre- 601 cipitate? [Ferric hydroxide, Fe 2 (OH) 6 or Fe(OH) 3 , simi- lar to iron rust.] What takes place when solution of ammonium sulphide is added to a ferrous solution ? Re- call the experiment made previously in studying the sul- phides. EXPLANATORY NOTE. The ferrous salts when in solution oxidize readily to ferric salts even by the action of the air, more quickly by the action of nitric acid. The color changes in consequence, and the precipitates with ammonium hydroxide and other reagents are different. Complete the following equations : Fe + HOI = FeCl 2 + Fe + H 2 S0 4 = FeS0 4 + FeS0 4 + HJSTOs = Fe 2 (S0 4 ) 3 + NO + FeCl 2 + NH 4 OH = Fe(OH) 2 + FeCl 3 + NH 4 HS = FeS + FeCl 3 + NH 4 OH = Fe(OH) 3 APPENDIX Weighing. Substances which corrode metal should be 1 weighed on glass, and general care should be taken to pro- tect the metallic parts of the balance from injury. Things when weighed should not be hot, but should be at the tem- perature of the room. During the more exact weighing, care should be taken that currents of air do not interfere with the swinging of the beam. The equilibrium should be tested while the beam is in motion by the equality of swing to either side of the zero point, $, on the scale. Generally the equilibrium should be tried with nothing in the pans before the actual weighing. In using balances of the type A or B (Fig. 5 or Fig. 6), it is best to put the ob- ject to be weighed on the left-hand pan, and the weights, or counterpoise, on the right. The weights should be placed in the pan for trial in regular succession, until one is found which is too heavy and one which is too light, and finally the exact counterpoise is found. It is well to read the weights first from the vacant places in the box, record the total, then read from the weights themselves as they are taken from the pan, the largest first. Care must be taken to return each weight to its proper place, to guard against the loss of any, and to leave the balance clean and in proper condition for the next weighing. The balance of the type A is for the heavier weighings, and will hardly respond to less than 0.2 or 0.3 of a gram. The balance B should indicate 0.01 of a gram. The equilib- rium is tested by pressing down the lever, Z, which raises 91 FIG. 5. Balance A. FIG. 6. Balance B. FIG. 7. Balance C. Weights. FIG. 8. FIG. 9. FIG. 10. APPENDIX 93 the beam so that it is free to swing. . The balance C has but one pan, and the weights are in the form of rings which slide along the beam without removal therefrom, but when not in use hang upont he peg, P. The largest ring gives grams in tens, the next smaller in units, the third in tenths, and the fourth in hundredths. So the total reading shown in the figure is 75.30 grams. Heating a test-tube, Fig. 11. A strip of paper is folded 2 two or three times upon itself and wrapped around the test-tube to serve as holder, or, better, a strip of asbestos is thus used. FIG. ll. FIG. 12. Pouring from a reagent bottle, Fig. 12. The stopper is 3 taken between the fingers of the right or left hand as pre- ferred, and withdrawn from the bottle. This is to avoid laying the stopper on the table. Inverting and shaking a test-tube, Fig. 13. 4 The mortar and pestle, Fig. 14. 5 FIG. 13. FIG. 14. 94 ELEMENTARY PRINCIPLES OF CHEMISTRY 6 The gas generator, Fig. 15. T, Thistle-tube ; D, deliv- ery tube. It is well to wet glass tubing when putting it through stoppers or connecting it with rubber hose. ^ To cut glass tubing. Make a scratch with a three- edged file where it is desired the tube shall break. Hold the tube in both hands, the thumbs brought together on the opposite side from the scratch. Pull the hands apart in the direction of the tube's axis, with only a very slight bending motion. If the tube is short or the glass thin, it is well to hold it in one or two folds of the towel. Fig. 16. 8 To smooth the sharp edge of a cut tube, hold the end in the gas flame until the glass softens. Fig. 17. 9 To bend glass tubing, hold it in the flame, preferably a wide one, twirling it between the thumb and fingers until the glass is thoroughly softened, then slowly bend to the desired angle. The bend should be smoothly curved, as seen in the figure. If the heating is insufficient, or the bending too hasty, the glass wrinkles, and the tube is likely to break. It is well to pass the tube through the flame once or twice after it is bent, so that the cooling may be slow. Fig. 18. 10 To draw out and close glass tubing, hold it in the flame until the glass is well softened, then draw apart slowly and to the desired extent, and finally heat it still more at the separating point and draw again. Fig. 19. 11 Heating a crucible, Fig. 20. The crucible, C, rests on the pipestem triangle, T, and this in turn on the iron ring of the stand, S. The tongs, X, are used to handle the crucible or its lid while hot. 12 Evaporation to dryness in the porcelain dish, D, Fig. 21. When there is danger of spattering, it is well to hold the burner in the hand, and to touch the tip of the small flame for a moment to the bottom of the dish. The pur- pose is to heat, but not enough to cause the formation of bubbles of vapor. Constant stirring with the rod, R, is important. FIG. 19. FIG. 15. FIG. 17. FIG. 20. FIG. 21. 96 ELEMENTARY PRINCIPLES OF CHEMISTRY 13 Filtration. Prepare the filter paper by folding it across two diameters at right angles to each other, then separating the quadrants so as to form a cone, d. The cone should be fitted into the funnel and then wet, so that the paper ad- heres to the glass. Fig. 22. Fig. 23. Filtering, pouring the liquid from the beaker, B, using the rod, R, upon the funnel, F', showing the stem of the funnel lying against the wall of the beaker, so that the liquid which runs through may not cause spattering by dropping into the middle of the beaker. 14 The graduated cylinder, Fig. 24, to use in measuring liquids; content, fifty cubic centimeters (c. c.). FIG. 24. FIG. 23. APPENDIX 97 The thermometer. The measurement of temperature 15A is made with the centigrade thermometer. The fixed points of this are the boiling point of water and the melting point of ice. The latter is made of the scale, and the former is made 100, the space between the two points being divided into one hundred equal parts, called degrees. In the Fahrenheit thermometer, which is more commonly used in matters other than scientific, these two points are made 32 and 212 respectively. The relation between the two scales is made clear in the diagram. Fig. 25. Thermometers are fragile and expensive, and must be used with care. J_op 32 BOILING POINT MELTING POINT FIG. 25. 98 ELEMENTARY PRINCIPLES OF CHEMISTRY 15B The barometer. Fig. 26. The long tube is completely filled with mercury, then inverted in the shallow dish of mercury. The atmospheric pressure on the free surface of the mercury sustains a column of varying length, the nor- mal being 760 millimeters. Fig. 27, the siphon form of barometer. \VACUUM FIG. 27. FIG. 26. TWENTIETH CENTURY TEXT-BOOKS. Uniform, I2mo NOW READY. Botanical Text-Books by JOHN MERLE COULTER, A. M., Ph. D., Head of Department of Botany, University of Chicago : Plant Relations. A First Book of Botany. Cloth, $1.10. Plant Structures. A Second Book of Botany. Cloth, $1.20. Plant Studies. An Elementary Botany. Cloth, $1.25. Plants. A Text-Book of Botany. Cloth, $i.So. Key to Some of the Common Flora. Limp cloth, 60 cents. A History of the American Nation. By ANDREW C. MCLAUGHLIN, A. M., LL. B. Cloth, $1.40. 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SEP 5 SEP 18 1936 APfl 6 1940 OCT 1940 AfK 19420 13 1043 FEB 9 1944 14 1946 "* LD 21-50m-l,'33 Vr> J / r>"-\r ID I OOvjO