OLAX*^^^^ 
 
 IN MEMORIAM 
 Edmund O'Neill 
 
A TEXT-BOOK OF 
 
 EXPERIMENTAL CHEMISTRY 
 
 (WITH DESCRIPTIVE NOTES) 
 
 LEE 
 
A TEXT-BOOK OF 
 
 EXPERIMENTAL CHEMISTRY 
 
 (WITH DESCRIPTIVE NOTES) 
 
 FOR STUDENTS OF 
 GENERAL INORGANIC CHEMISTRY 
 
 BY 
 
 EDWIN LEE 
 
 PROFESSOR OF CHEMISTRY IN ALLEGHENY COLLEGE 
 
 WITH 57 ILLUSTRATIONS 
 
 PHILADELPHIA 
 P. BLAKISTON'S SON & CO, 
 
 1012 WALNUT STREET 
 1908 
 
COPYRIGHT, 1908, BY P. BLAKISTON'S SON & CO. 
 IN MEMORIAM 
 
 Printed by 
 
 The Maple Press 
 
 York Pa. 
 
PREFACE. 
 
 WHILE no particular claim to originality is made for this text-book, as 
 many of the experiments have been described previously, yet the writer 
 believes that the book will be found to be something more than a mere 
 compilation. It grew originally out of a personal demand for a text- 
 book which would embody: (a) a clear, accurate and comprehensive 
 presentation of the fundamentals of the science; (b) specific directions 
 for laboratory work, coupled with such questions as lead the student to 
 observe, compare and generalize, and w r ould therefore provide a method 
 for the scientific development of the principles under discussion; (c) a 
 sufficient amount of discussion and application of the principles involved 
 in the experiments to foster the interest and to direct the observa- 
 tions that energy may not be spent indiscriminately, and (d) those physico- 
 chemical generalizations which are essential to the explanation of much 
 of the phenomena of inorganic chemistry. 
 
 This book represents an endeavor to meet these requirements. It 
 is not intended that it shall take the place of a large descriptive work or 
 the instruction of the teacher; on the contrary, it is designed to provide, 
 primarily, an experimental course in general chemistry, and by the use of 
 "descriptive notes" and questions vitally relate it to the lecture-room 
 work. It is scarcely necessary to emphazise the importance of laboratory 
 work as being essential to a thorough comprehension of the subject; 
 but this same work has a very doubtful value unless it is carefully directed 
 and correlated with the lecture and text-book. It must not be merely a 
 mechanical part of the course. The student must see that his laboratory 
 work is but a means to an end that lectures and experiments are mu- 
 tually helpful. Very frequently the laboratory work is taught too much 
 apart from the course. Beginners often complain, and more frequently 
 conduct their work as if no relation existed between lectures and labora- 
 tory work. It is contended, therefore, that the laboratory manual should 
 provide something to make obvious this relationship and to assist in the 
 fusion of the two. It is hoped that the " notes" appended to many of the 
 experiments, and which embody discussions, applications, formulae, etc., 
 will form the connecting link between the work of the laboratory and 
 
VI PREFACE. 
 
 the work of the lecture-room. In addition to the foregoing, sufficient 
 amounts of descriptive, theoretical and physical chemistry have been 
 incorporated to present in outline, at least, the essentials of general 
 chemistry. It is presumed, however, that the laboratory work will be 
 accompanied by a full course of lectures. 
 
 The classification is according to the natural families of the periodic 
 system. 
 
 The fundamental concepts of the science are first built up, explained and 
 illustrated, and then applied persistently. 
 
 Neither the inductive nor the deductive method has been adhered 
 to with "faddish" tenacity. Experience has taught the author that a 
 combination of the two methods can be relied upon to yield satisfactory 
 results. The inductive method possesses undoubted virtues, yet experi- 
 enced teachers cannot fail to appreciate the pertinency and significance 
 of Faraday's remark, "What are we to look for, Tyndall," as the 
 latter was about to perform an experiment. A too vigorous application 
 of even this method usually results in a waste of energy. There is 
 danger in any extreme. The notes and questions in the book are arranged 
 with a view to imparting definiteness of purpose and aim to the labora- 
 tory work. 
 
 Most of the exercises have been used in the type-written form for 
 a number of years. They cannot be regarded as "impracticable and 
 beyond the capabilities of the average student." An effort has been 
 made to lighten the labor of the teacher by making the directions specific. 
 Further, a series of graduated questions have been incorporated among 
 these directions in order to asisst the student in correlating and generaliz- 
 ing. This also should save the teacher's time and energy. The rigid- 
 ness imparted by the use of specific directions is overcome by the intro- 
 duction of abundant material to meet varied conditions a much larger 
 number of experiments being inserted than will likely be used either in 
 lecture-room or laboratory. This enables the teacher to make a selection. 
 
 No hesitancy has been manifested in introducing physico-chemical 
 generalizations wherever they have seemed necessary to rational correla- 
 tion and explanation of facts. The importance of these generalizations 
 in connection with the teaching of general chemistry is now generally 
 recognized, yet the writer confesses that he knows of no text-book on 
 "experimental chemistry" which has been brought abreast of the times 
 by appropriating and incorporating the results of the labors of the physical 
 chemist. It may be true that physical chemistry cannot and should not be 
 
PREFACE. VI] 
 
 taught in the first year of chemistry, but in the opinion of the author and 
 at least one other, "many of these generalizations which have been brought 
 in by this new physical chemistry, and which affect fundamentally the 
 whole science of chemistry, can and should be introduced into general 
 chemistry." The author w r ould not leave the impression that the older 
 generalizations and methods of developing the subject have been aban- 
 doned; on the contrary, they have been retained and an effort made tc 
 rationalize them by supplementing them with more recent data taken 
 from the domain of physical chemistry. The book represents, therefore, 
 an effort to fuse modern views and recent advances with those older views 
 which have stood the test, and to explain chemical phenomena in a man- 
 ner that is in accord with modern chemical thought. 
 
 The role of energy in chemical reaction has been given unusual promi- 
 nence. The elements of thermochemistry have been presented. More 
 space has been allotted to "solutions," "the ion theory," "chemical 
 equilibrium" and the "mass law" than is usually given in books of this 
 character. It is believed, however, that the results obtained warrant 
 this. A brief comprehensive statement of the historical development 
 of "the electrolytic dissociation theory" has been given that the student 
 may become familiar with the story of the gradual development of at 
 least one chemical theory. The author has observed that students take a 
 keen delight in stating just how much was contributed to the develop- 
 ment of a chemical theory by this or that chemist. With a similar purpose 
 in view, the writer has quoted references verbatim, rather than record the 
 gist of them. 
 
 Chapter VII presents in outline such subjects as the kinetic-molecular 
 hypothesis, atomic theory, valence, formulae, equations and units. This 
 gives flexibility to the course by allowing the teacher to develop the 
 several subjects in a manner best suited to the methods in use. 
 
 The consulting of easily-found references is demanded frequently. 
 The London Chemical News says in this connection: "The habit of 
 looking up whatever needs greater elucidation augurs well for good 
 culture and education." If possible, many of the books mentioned in 
 the reference list should be made accessible to the students. 
 
 The author's supreme purpose has been to present a system of chemistry 
 rather than an assemblage of chemical facts, and to build a book which 
 will aid the student in cultivating correct habits of thought. 
 
 The author acknowledges with pleasure the help and inspiration of 
 those masterful teachers with whom he has been permitted to work. 
 
Vlll PREFACE. 
 
 To his fellow teachers who have been generous with valuable sugges- 
 tions, he acknowledges his indebtedness. The writer desires especially 
 to thank Professor Arthur B. Lamb, of New York University, and 
 Dr. J. H. Mathews of the University of Wisconsin for helpful criticisms. 
 
 For courtesies received, the writer hereby extends his thanks to 
 Professor T. W. Richards of Harvard University, Professor A. A. Noyes 
 and Dr. G. N. Lewis of Massachusetts Institute of Technology, Professor 
 Harry C. Jones and Dr. H. N. Morse of Johns Hopkins University, Professor 
 T. L. R. Morgan of Columbia University, Professor Alexander Smith of 
 the University of Chicago, and Professor F. A. Gooch of Yale University. 
 
 I am indebted to my publishers, P. Blakiston's Son & Company, 
 for the use of certain illustrations from several of their publications 
 and desire to express my appreciation of their courtesy. 
 
 The author will be grateful to those who report errors and ambiguity. 
 
 ALLEGHENY COLLEGE, EDWIN LEE. 
 
 Meadville, Penna., June, 1908. 
 
CONTENTS. 
 
 INTRODUCTORY. 
 
 PAGE 
 
 CHAPTER I. 
 FUNDAMENTAL CONCEPTS 17 
 
 Relationships and Definitions. Entity Defined. Physics versus 
 Chemistry. Note on Matter. "System" Defined. Specimens or 
 Bodies. Substances. Homogeneity and Heterogeneity. Proper- 
 ties. Identity. Experiments. Note on Energy. Forms of Energy. 
 First Law of Energy. Transformation of Energy. Experiments . 
 Relation of Energy to Properties of Matter. Phase. Relation of 
 Energy to Phase of a Substance. Experiments. 
 
 CHAPTER II. 
 
 CHARACTERISTICS OF PHYSICAL AND CHEMICAL CHANGES 26 
 
 Note. Definitions. Physical and Chemical Changes. Electrolysis 
 of Water. Change of Identity. Analysis. Effect of Heat on Mer- 
 curic Oxide. Effect of Heating Copper in Air. Synthesis. Solution 
 Facilitates Chemical Action. Note on Surface Phenomena. Relation 
 of Energy to Chemical Change. Exercises. 
 
 CHAPTER III. 
 
 ELEMENTARY SUBSTANCES, MIXTURES AND COMPOUNDS 31 
 
 Note. Second Law of Matter. Examination of the Elementary Sub- 
 stances. Energy Content Affects Properties of Elementary Sub- 
 stances. Mixtures Heterogeneous and Homogeneous. Note. 
 Mixtures and Compounds. Definitions. Interaction of Compounds. 
 Metathesis. Reaction. Reagents. "Factors." "Products." 
 
 CHAPTER IV. 
 
 NOTE ON THE ENERGETICS OF CHEMISTRY 34 
 
 Theory. First Law of Energy. "Matter Content." "Energy 
 Content." Internal Energy, Cohesion Energy, Disgregation Energy, 
 Chemical Energy. Free and Bound Energy. LeChatelier's Theorem. 
 Application of Theory. Red and Yellow Phosphorus. Allotropy. 
 Thermochemistry. Endothermic and Exothermic Reactions. First 
 Law of Thermochemistry. Law of Hess. Heat of Reaction versus 
 Heat of Formation. 
 
 CHAPTER V. 
 
 SCIENCE, ITS METHODS OF DEVELOPMENT, CLASSIFICATION 40 
 
 Definitions of Phenomenon, Experiment, Fact, Law, Rule. Induc- 
 tion, Deduction. Hypothesis, Theory, Prediction. Scientific versus 
 Systematic Methods. Physical Sciences and Biological Sciences. 
 Abstract and Concrete Sciences. 
 
x CONTENTS. 
 
 CHAPTER VI. 
 
 PAGE 
 
 QUANTITATIVE RELATIONSHIPS. LAWS AND CHEMICAL EQUIVALENTS ... 43 
 
 N ote- Conservation of Mass in a Limited System. Analysis of 
 Water. First Law of Matter. Combination in Definite Propor- 
 tions by Weight. Synthesis of Copper Oxide, Magnesium Oxide, 
 Iron Oxide. Law of Definite Proportions. Constancy of Composi- 
 tion. Combination of Two Substances in Varying Proportions by 
 Weight. Law of Multiple Proportions. Combining Volumes. 
 Synthesis of Water from Hydrogen and Oxygen Gases. Syntheses of 
 Steam. The Law of Gay-Lussac. Law of Reciprocal Proportions. 
 Chemical Equivalents. Equivalent Weight of Zinc by Displacement 
 of Hydrogen. 
 
 CHAPTER VII. 
 
 OUTLINES, THEORIES, FORMULAE, VALENCE, EQUATIONS, UNITS 50 
 
 Structure of Matter. Molecule. Kinetic-Molecular Hypothesis, and 
 Its Application. Critical Phenomena. Combining Volumes. 
 Atom. Atomic Theory and Its Application. Electronic Theory of 
 Matter. Nomenclature of Atomic Theory. Simple and Compound 
 Molecules. Molecular Weight. Mass and Energy Equations. 
 Note. Valence. Multiple Valence. Kinds of Formulae. Units of 
 Measurement. Stoichiometry. Methods for the Determination of 
 Molecular and Atomic Weights. References. Problems. 
 
 CHAPTER VIII. 
 
 OXYGEN 54 
 
 Introductory Note on Occurrence and Properties of Oxygen. Sources 
 of Oxygen. Laboratory Sources of Oxygen. Oxidation of Carbon, 
 Phosphorus, Sulphur, Iron. Oxides. "Ous" and "ic" Endings. 
 Kindling Temperature. Note on Heat of Combustion. Energy 
 Equations. Catalytic Action of Manganese Dioxide. Note on 
 Catalyzers. Preparation and Properties of Ozone. Note on Energy 
 of Oxygen. Oxygen and Ozone. Determination of the Density 
 Relations of Determination of the Volume of Oxygen Liberated from 
 a Given Weight of Potassium Chlorate. Percentage of Oxygen in 
 the Air. Dalton's Law of Partial Pressures. Problems. 
 
 CHAPTER IX. 
 
 HYDROGEN 64 
 
 Note. Methods for Preparing Hydrogen. Displacement of Hydro- 
 gen from Acids by Metals. Note on Change of Energy in a System. 
 Properties of Hydrogen. Effusion of Hydrogen. Graham's Law of 
 Effusion. Occlusion of Hydrogen. Catalytic Action of Spongy 
 Platinum. Note on Catalyzers and Surface Phenomena. Reduction 
 of Cupric Oxide. The Nascent State. Nascent Hydrogen. Prepara- 
 tion and Properties of Hydrogen Dioxide. Hydrogen Equivalent of 
 Zinc. Outline of Hydrogen. TJaeHttochernical Deportment of Hydro- 
 gen and Oxygen in Formation of Water and Hydrogen Dioxide. 
 Problems. 
 
CONTENTS. XI 
 
 CHAPTER X. 
 
 PAGE 
 
 WATER 75 
 
 Note. Composition and Properties of Water. Purity. Temper- 
 ature. Density Graph of Water. Hydrolysis. Chemical Union with 
 Oxides. Water of Hydration. Vapor Tension of Substances, Efflor- 
 escence and Deliquescence. Determination of Water of Crystalliza- 
 tion. Outline of Properties of Water. Note on Polymerization and 
 Hydrolysis. 
 
 CHAPTER XI. 
 
 SOLUTIONS 82 
 
 Note and Definitions. Kinds of Solutions. Solubility of Solids. 
 Characteristics of a Solution. Note on Colloidal Solutions. Surface 
 and Diffusion Phenomena. Solution Tension. Note on Equilibria in 
 Solutions. Osmotic Pressure. Pfeffer's Data. Morse and Frazer's 
 Results. Van't Hoff's Law. Effect of Temperature on Solubility of 
 Solids. Supersaturated Solutions. Thermal Phenomena. Relative 
 Solubility of Solids. Terminology of Solutions. Table of Solubility 
 of Solids. Note on Heat of Solution. Table, Heat of Solution of 
 Solids, Liquids and Gases. Effect of Pressure on Solubility. Solu- 
 bility of Pairs of Liquids. Relative and Mutual Solubility. Effect of 
 Temperature. Critical Solution Temperature. Distribution of 
 Solute Between Two Immiscible Solvents. Law of Distribution. 
 Solubility of Gases. Pressure of Gas Affects Solubility. Henry's 
 Law. Effect of Solvent on Solubility. Effect of Temperature on 
 Solubility of Gases. Elevation of the Boiling Point. Depression of 
 the Freezing Point. Notes. Raoult's Laws. Lowering of the Solu- 
 tion Tension. Note. Problems. 
 
 CHAPTER XII. 
 
 ACIDS, BASES AND SALTS 104 
 
 Chlorine. Preparation and Experimental Study of Its Properties. 
 Note. Hydrogen Chloride Gas. Acids. General and Specific 
 Properties of Acids. Anhydrides. Acid Radicals. Classification 
 of Acids: Sodium. Its Properties. Note. Sodium Hydroxide. 
 Bases General and Specific Properties. Classification of Bases. 
 Interaction of Acids and Bases. Neutralization.: Salts. Classifica- 
 tion and Nomenclature of Salts. Metal and Non-Metal Defined. 
 Problems. 
 
 CHAPTER XIII. 
 
 CHEMICAL EQUILIBRIUM, REVERSIBLE REACTIONS, MASS LAW, DISSOCIATION, 1 19 
 
 Note on Chemical Equilibrium and Law of Concentration Effect. 
 
 Mathematical Form of Mass Law. Affinity Constant. Reactions 
 
 are Incom- 
 a Gas. 
 
 Which are Approximately Complete. Reactions which ; 
 plete. Equilibrium. Berthollet's Laws. Dissociation of 
 
 CHAPTER XIV. 
 
 NOTE ON THE MODERN THEORY OF SOLUTION 124 
 
 Historical Development of the Electrolytic Dissociation Theory. 
 Statement of the Theory. Le Blanc, Nernst, Thomson and Beilby, 
 quoted in regard to the probable source of the electric charges on the 
 ions. Application of Theory. Electrolysis. Electro-Chemical Equiv- 
 alent. Names of Ions. Electrolytes, Half-Electrolytes and Non- 
 
Xll CONTENTS. 
 
 PAGE 
 
 Electrolytes. Note. lonization in Solution. Effect of the Nature of 
 the Solvent and the Solute on lonization. Effect of Dilution. 
 Ostwald's Law of Dilution. Dissociation Constant. Color of Ions. 
 Chemical Conduct of Ions. "Strength" of Acids and Bases. 
 Tables. Thermochemical Support of the Dissociation Theory. 
 Change in Volume Support. Salts. Compound Salts versus Complex 
 Salts. Hydrolysis. Ionic Equilibrium. "Ion Product." "Solu- 
 bility Product." Solution Tension. Relation of Chemical Energy 
 and Electrical Energy. Note. E.M.F. Is a Measure of Chemical 
 Energy When the Process is Reversible. Table of Single Potentials. 
 Heat of lonization. Table. 
 
 THE NON-METALS OR ACID-FORMING ELEMENTS. 
 CHAPTER XV. 
 
 THE CHLORINE FAMILY 152 
 
 Introductory Note. Preparation and Properties of Fluorine. Hydro- 
 gen Fluoride. Silicon Tetrafluoride. Note: Chlorine, Preparation 
 and Properties. Hydrogen Chloride. Chlorides. Oxygen Deriva- 
 tives of Chlorine. Chlorates and Perchlorates : Bromine, Prepara- 
 tion and Properties. Substituting Power of Chlorine as Compared 
 with Bromine. Hydrogen Derivatives. Bromides. Oxygen Deriva- 
 tives. Bromates, etc.: Iodine, Preparation and Properties. Tinc- 
 tures. Substituting Power of Chlorine and Bromine as Compared with 
 Iodine. Liberation of Iodine by Oxidizing Agents. Displacement 
 of Sulphur from Hydrogen Sulphide by Chlorine, Bromine or Iodine. 
 Hydrogen Derivatives of Iodine. Iodides. Oxygen Derivatives. 
 lodates. etc.: Periods of Induction in Chemical Reactions. Resem- 
 blance of the Properties of Corresponding Compounds of the Halogens. 
 Detection of Fluorides, Chlorides, Bromides, and Iodides, together 
 in a Solution. Comparative Table. Note. 
 
 CHAPTER XVI. 
 
 CLASSIFICATION OF THE ELEMENTS 167 
 
 Note. Prout's Hypothesis. Dobereiner's Triads. Newland's Law 
 of Octaves. Mendeleef's Original Table. Lothar Meyer's Natural 
 Classification. Periodic Arrangement. Bayley's Table. Atomic 
 Volume Table. Table of Atomic Volume and Specific Gravity. 
 Exercises. 
 
 CHAPTER XVII. 
 
 GROUP VI. FAMILY M. THE OXYGEN FAMILY 176 
 
 Note. Oxygen. Hydrogen Derivatives and Their Thermal Con- 
 duct. Allotropy. Ozone : Sulphur. Allotropy. Note. Sulphides. 
 A Test for Sulphides. Oxygen Derivatives of Sulphur. Sul- 
 phuric Acid. Manufacture of Sulphuric Acid. Soluble and 
 Insoluble Sulphates. Contraction in Volume When Sulphuric Acid 
 and Water are Mixed. Table: Selenium: Tellurium: Compara- 
 tive Table. Problems. 
 
 CHAPTER XVIII. 
 
 NITROGEN AND THE ATMOSPHERE. THE HELIUM FAMILY 192 
 
 Nitrogen. Preparation and Properties. Determination of the Con- 
 stituents of the Atmosphere. Weight of a Liter of Air: The Helium 
 Family. Note. Tables. Problems. 
 
CONTENTS. Xlll 
 
 CHAPTER XIX. 
 
 PAGE 
 
 GROUP V. FAMILY M. NITROGEN FAMILY 199 
 
 Note on Relationships. Nitrogen, Preparation and Properties. 
 Hydrogen Derivatives. Ammonia. Ammonium Salts. Nitrogen and 
 the Halogens. Preparation of an Endothermic Compound, Nitrogen 
 Tri-iodide. Composition of Ammonia Gas. Oxides of Nitrogen are 
 Endothermic Compounds. Note on Heat of Formation. Nitric 
 Acid. Note and Equations Relative to the Conduct of Nitric Acid. 
 Aqua Regia. Nitrates: Phosphorus. Allotropy. Hydrogen Deri- 
 vatives. Phosphorus and the Halogens. Oxygen Derivatives. 
 Phosphoric Acid. Tests to Distinguish the Phosphates: Arsenic. 
 Properties. Derivatives. Tests for Arsenic: Antimony. Deriva- 
 tives. Hydrolysis of Antimony Trichloride. Sulphides of Antimony: 
 Bismuth. Properties. Alloys. Hydrolysis of Bismuth Nitrate. 
 Comparative Table. Problems. 
 
 CHAPTER XX. 
 
 CARBON FAMILY 226 
 
 Primary Group. Carbon. Allotropy. Carbon and Oxygen. 
 Energy Equations. Carbonates. Application of the Law of Hess. 
 Carbon and Hydrogen (Hydrocarbons). Acetylene. Coal Gas. 
 Carbon Disulphide. Carbon and the Halogens. Carbon and Nitro- 
 gen. Cyanides. A study of Flames. Temperature of Flame of 
 Bunsen Burner: Silicon. Note. Silic Acid. Dialysis. Fluo- 
 silicic Acid. A Test for Silicates. Note. Comparative Table. 
 Problems. 
 
 CHAPTER XXI. 
 
 SOME COMMON CARBON COMPOUNDS 246 
 
 Introductory Note. Hydrocarbons. Alcohols. Phenols. Ethers. 
 Aldehydes. Acids. Ketones. Halides of Carbon. Esters. Ammo- 
 nia. Derivatives. Organo-Mineral Compounds. Carbohydrates. 
 Glucosides. Albuminoids. Isomerism. Flashing-point of Kerosene. 
 Tests for Ethyl Alcohol. Glycerin. Soaps. Carbolic Acid. 
 Glacial Acetic Acid. Tartaric Acid. Oxalic Acid. Chloroform. 
 lodoform. Butter. Fatty or Fixed Oils. Essential or Volatile Oils 
 Sugars. Starches. Gun-Cotton or Smokeless Powder. Paper. 
 
 CHAPTER XXII. 
 
 ALUMINUM FAMILY 271 
 
 Note. Boron. Properties. Borax. Solubility of Metallic Oxides in 
 Borax. Borax. Bead Tests. Table. Detection of Boron. Com- 
 parative Table. 
 
 THE METALS OR BASE-FORMING ELEMENTS. 
 CHAPTER XXIII. 
 
 INTRODUCTORY NOTE 2 7 6 
 
 Physical Properties of the Metals. Chemical Properties of the Metals. 
 Occurrence of the Metals in Nature. Table. Classification of the 
 Metals. 
 
XIV CONTENTS. 
 
 CHAPTER XXIV. 
 
 PAGE 
 
 GROUP I. FAMILY M. ALKALI METALS 280 
 
 Introductory Note. Lithium. Flame Color of Lithium Compounds : 
 Sodium. Experiments. Tests. Alkalimetry and Acidimetry: Po- 
 tassium. Experiments. Tests. Gunpowder: Ammonium. Experi- 
 ments. Detection of the Alkali Metals in a Mixture: Rubidium: 
 Caesium: Problems. 
 
 CHAPTER XXV. 
 
 GROUP I. FAMILY M. THE COPPER FAMILY 292 
 
 Introductory Note. Copper. Experiments. Scheme of lonization 
 for Double Salts and Complex Salts. Note: Silver. Experiments. 
 Photography: Gold. Experiments. "Purple of Cassius:" Prob- 
 lems. 
 
 CHAPTER XXVI. 
 
 GROUP II. FAMILY M. ALKALINE EARTH METALS 306 
 
 Introductory Note. Relative Solubility of Salts. Glucinum: 
 Magnesium. Experiments : Calcium. Experiments. Quicklime. 
 Plaster of Paris. "Permanent Hardness of Water:" Strontium. 
 Experiments: Barium. Experiments: Detection of the Alkaline 
 Earth Metals in a Mixture: Radium: Exercises. Problems. 
 
 CHAPTER XXVII. 
 SECONDARY FAMILY OF THE ALKALINE EARTHS 316 
 
 Introductory Note. Zinc. Comparative Table. Experiments, 
 An Experimental Study of Ionic Equilibrium and "Concentration 
 Effect:" Cadmium. Formation of Salts and Analytical Reactions: 
 Mercury. Properties. Formation of Salts and Analytical Reactions: 
 Problems. 
 
 CHAPTER XXVIII. 
 
 THE ELEMENTS OF GROUP III 326 
 
 Family M Scandium, Yttrium, Lanthanum, Ytterbium. Note. 
 
 Absorption and Emission Spectra. Eka-boron and Eka-Aluminum of 
 
 Mendeleeff. 
 
 Family m Note. Boron. Aluminum. Experiments. "Thermit." 
 
 Ruby. Sapphire. Emery. Clay. Classification of Alums: 
 
 Thallium. Properties. Salts of Thallium: Problems, 
 
 CHAPTER XXIX. 
 
 THE ELEMENTS OF GROUP IV 333 
 
 Family M Titanium, Zirconium, Cerium, Thorium. Note. 
 Family m Note. Tin-plate. Type-metal. Shot-metal. German- 
 ium: Tin. Experiments. Determination of the Equivalent Weight 
 of Tin. Analytical Reactions: Lead. "Lead Tree." Formation 
 and Preparation of Salts. Chemical Principles Involved in the Manu- 
 facture of White Lead : Problems. 
 
 CHAPTER XXX. 
 
 ELEMENTS OF GROUP V 343 
 
 Family M Vanadium, Columbium, Tantalum. Note. 
 
CONTENTS. XV 
 
 CHAPTER XXXI. 
 
 PAGE 
 ELEMENTS or GROUP VI 344 
 
 Family M Chromium. Note. Experiments. Chromyl Chloride, 
 a Test for Chlorides. Analytical Reactions: Molybdenum. Ammo- 
 nium Phosphomolybdate : Tungsten. Tungsten Steel: Uranium. 
 Pitchblende. Radio-activity : Problems. 
 
 CHAPTER XXXII. 
 
 ELEMENTS OF GROUP VII 351 
 
 Family M Manganese. Note. Experiments. Equations. Analyt- 
 ical Reactions: Problems. 
 
 CHAPTER XXXIII. 
 
 TRANSITION ELEMENTS 355 
 
 First Long Period. Note: Irori. Note. Application of Iron to 
 Industrial Arts. "Passive State." Corrosion of Iron. Experiments. 
 Reactions of Ferrous and Ferric Salts. Ferrocyanides and Ferricy- 
 anides. Oxidation and Reduction of Iron Salts: Cobalt. Note. 
 Experiments. Reactions : Nickel. Note. Electro-plating. Alloys : 
 Problems. 
 
 CHAPTER XXXIV. 
 
 TRANSITION ELEMENTS 366 
 
 Second Long Period. Note: Ruthenium. Note: Rhodium. 
 Note: Palladium. Note. Occlusion "Solid Solutions." Fourth 
 Long Period. Note: Osmium. Note: Iridium. Note: Platinum. 
 Note. Platinum Black. Experiments. Platinum as a Catalyzing 
 Agent. Reactions. 
 
 CHAPTER XXXV. 
 
 RELATIONS WITHIN THE GROUPS OF THE PERIODIC CLASSIFICATION .... 370 
 Grouping of the Metals (Cations) for Purposes of Analysis; Grouping 
 of Non-Metallic Radicals (Anions) for Purposes of Analysis. 
 
 APPENDIX I 373 
 
 PRELIMINARY EXERCISES 
 
 The Bunsen Burner. The Blast-Lamp. Manipulation of Glass. 
 Perforation of Stoppers. Treatment of Rubber Corks and Tubing. 
 The Wash Bottle. Measuring Instruments. Calibration by Weigh- 
 ing. Problems 
 
 APPENDIX II (Tables) 381 
 
 REFERENCE BOOKS 415 
 
 LABORATORY INSTRUCTIONS xx 
 
 INDEX ... 419 
 
LIST OF TABLES. 
 
 PAGE. 
 
 Osmotic Pressure of Cane Sugar 86 
 
 Osmotic Pressure and Molecular Weights 87 
 
 Solubility of Solids 91 
 
 Heat of Solution of Solids 92 
 
 Heat of Solution of Liquids 92 
 
 Heat of Solution of Gases 93 
 
 Solubility of Gases 96-97 
 
 Constants (Elevation of the Boiling Point) 99 
 
 Raoult's Tables (Depression of the Freezing Point) 101 
 
 Names of Ions 135 
 
 Velocity of Ions 135 
 
 Heat of Neutralization 143 
 
 Acids, Bases, Salts. Per Cent. Ionized 145 
 
 Order of Solution Tension of the Metals 147 
 
 Single Potentials 150 
 
 Heat of lonization 151 
 
 Comparative Table. The Halogens 165 
 
 List of Non-Metals 167 
 
 Triads of Dobereiner 168 
 
 Newland's Octaves 168 
 
 Mendeleef's Original Table 169 
 
 Lothar Meyer's Original Table 170 
 
 Periodic arrangement: 
 
 Table 1 171 
 
 Table II 172 
 
 Bayley's Table 173 
 
 Atomic Volume, Specific Heat, Valence 1 74 
 
 Atomic Volume Specific Gravity 175 
 
 Comparative Table. The Oxygen Family 191 
 
 Comparative Table. The Helium Family 197 
 
 Comparative Table. The Nitrogen Family 225 
 
 Temperature of Bunsen Burner Flame 240 
 
 Comparative Table. Carbon and Silicon 244 
 
 Colors of Borax Beads 274 
 
 Comparative Table. Boron and Aluminum 275 
 
 Comparative Table. Zinc, Cadmium and Mercury 317 
 
 Comparative Table. Germanium, Tin and Lead 337 
 
 xvii 
 
XViii LIST OF TABLES. 
 
 PAGE 
 
 Grouping of the Metals (Cations) for Purposes of Analysis 370 
 
 Grouping of the Non-Metallic Radicals (Anions) for Purposes of Analysis . . . .371 
 
 Metric Measures with English Equivalents 381 
 
 English Measures with Metric Equivalents 382 
 
 Conversion of Thermometric Readings 383 
 
 Corrections for Mercury-in-Glass Thermometers 383 
 
 Density of Water 384 
 
 Volume of One Gram of Water at Various Temperatures 385 
 
 Density of Mercury 385 
 
 Barometric Corrections 386 
 
 Reduction to Vacuum of Weighings Made in Air 387 
 
 Correction Factors for Calibrating Glass Vessels 388 
 
 Vapor Tension of Water (o-35) 389 
 
 Vapor Tension of Water (o-ioo) 390 
 
 Vapor Pressure of Mercury 391 
 
 Boiling Temperature (t) of Water at Barometric Pressure (b) 392 
 
 Boiling Temperature (t) of Water at Pressure of (a) Atmospheres 392 
 
 Physical Constants 393 
 
 Composition of the Air by Volume 394 
 
 Diffusion of Gases 394 
 
 Specific Heat Atomic Heat 395 
 
 Heat of Formation 
 
 (a) Chlorides, Bromides, Iodides 397 
 
 (b) Sulphates, Nitrates, Carbonates 398 
 
 (c) Oxides, Hydroxides, Sulphides 399 
 
 Solubilities 400 
 
 Solubilities of Salts (Quant.) 401 
 
 Composition of Some of the Important Alloys 402 
 
 Scale of Hardness 404 
 
 Color Scale of Temperature 404 
 
 Indicators 405 
 
 Soap-Bubble Solutions 405 
 
 Graduated Solutions: 
 
 (a) Fehling's Solution 406 
 
 (b) Nessler's Solution 406 
 
 Percentage and Specific Gravity of Solutions at 15 C.: 
 
 (a) Sulphuric Acid 407 
 
 (b) Hydrochloric Acid 408 
 
 (c) Nitric Acid 408 
 
 (d) Ammonium Hydroxide 409 
 
 Proportion by Weight of Absolute Alcohol 410 
 
 Proportion by Volume of Absolute Alcohol 411 
 
 List of Elements, Valences, Discoverer, etc 412-414 
 
LABORATORY INSTRUCTIONS. 
 
 After receiving assignment of desk, procure key from rack. Check your appara- 
 tus by comparison with list found in drawer. Report any differences at once to the 
 assistant. 
 
 Read the "Regulations" posted in various places in the laboratory. Always read, 
 in entirety, the directions for performing an experiment before assembling apparatus. 
 
 It is well to provide yourself with apron or blouse to protect clothing while at work , 
 also white soft cloth to be used for wiping apparatus. A sponge is convenient. 
 
 Procure a note-book (Instructions). Make a neat, permanent and true report of 
 each experiment immediately after its performance under the following heads: 
 I. Object of Experiment. 
 II. Manipulation. (What you did.) 
 
 III. Observed Phenomena. 
 
 IV. Conclusions or Results. 
 
 V. Give equations if requested. 
 VI. Answer questions. 
 
 VII. Errata (any mistakes should be noted under this head). 
 VIII. A diagram of apparatus frequently facilitates the interpretation of an 
 experiment. 
 
 Students must work independently, both as to manipulation and records, unless 
 otherwise specified. 
 
 When weighing is necessary, use rough balances unless experiment is marked 
 "Quant." 
 
 Do not carry the bottles containing the various substances from the shelves to your 
 desk. Keep the bottles in order. Use a test-tube or watch glass or a piece of paper 
 to transfer substances. Do not return to the bottles unused portions of substances 
 unless you have secured permission from the instructor. If quantity to be used is 
 not specified use 2 cm.3 or 3 cm.3 
 
 Do not throw anything away until quite sure that you are through with it. Throw 
 waste liquids into sink; other waste material into waste jars. 
 
 When replacing apparatus use order-sheet (pink). 
 
 T. O. means that apparatus for experiment is to be secured temporarily from the 
 instructor; i.e., it is to be returned to him after performance of experiment. 
 
 L. T. suggests that experiment can be performed to advantage on lecture -table. 
 
 Examine your desk on entering the laboratory. If anything has been disturbed or is 
 missing, report same at once to assistant. 
 
 Before leaving the laboratory, place your desk in order. Cleanliness is absolutely 
 necessary. Failure to observe this contributes to failure. After observing these instruc- 
 tions, lock your desk and place key on rack. See that gas and water are turned off. 
 
 NOTE. The student will be graded on his "laboratory deportment"; i. e., the 
 persistency with which the foregoing instruct ons are adhered to. 
 
 XX 
 
A TEXT-BOOK 
 OF 
 
 EXPERIMENTAL CHEMISTRY. 
 
 CHAPTER I. 
 FUNDAMENTAL CONCEPTS: 
 
 RELATIONSHIPS AND DEFINITIONS. 
 
 Through the medium of our senses we are constantly receiving sensa- 
 tions which we interpret objectively, i.e., we locate the cause of a sensation 
 in a particular portion of space. As the result of our experience we assume 
 that the physical universe has an objective existence, and that our ac- 
 quaintance with it depends solely upon; otic senses. { ' 
 
 Men have given the name "thing." or "entity*' to the'cause of a sen- 
 sation to that which has the objective- exisfe'uojr: '-^he' entities with 
 which the scientist is particularly concerned are Matter and Energy. 
 Time, Temperature, Space, Velocity, etc., are not things.* 
 
 It will be readily recalled from the student's experience in the study of 
 physics that the two classes of things with which the physicist has to do 
 are those previously enumerated. The same is true of the chemist. 
 
 The fundamental difference between Physics and Chemistry rests 
 upon the relative amount of emphasis laid upon either Energy or Matter. 
 
 Physics is essentially the science of Energy, and aside from a compara- 
 tively brief discussion of the properties of Matter, considers the latter 
 only as it is associated with or becomes "the vehicle of Energy." 
 
 Chemistry has been termed the science of matter. This is due to the 
 fact that heretofore but little time has been devoted to energy considera- 
 tions. Recent years, however, have witnessed a marked change in the 
 methods employed in the study of Chemistry. More emphasis is being 
 placed upon the role of energy in this science. With this change, the 
 artificial line of demarcation, separating Physics and Chemistry, has 
 practically disappeared. 
 
 MATTER. 
 
 If we consider the evidence furnished by scientific investigation, it is 
 difficult to give a satisfactory and final definition of matter. It is better 
 described by its properties, although it is somewhat evasively defined as 
 anything which occupies space and appeals to the senses. From a 
 
 *NOTE: A few scientists maintain that there is no such thing as "matter," that it 
 is but an energy manifestation; others hold that neither matter nor energy is an ob- 
 jective reality, but merely a product of thought. 
 
 17 
 
!8 EXPERIMENTAL CHEMISTRY. 
 
 chemical point of view, it has been described as anything which possesses 
 weight. Science seems unable as yet to predicate what matter is. Ex- 
 periment has revealed that "the total mass (quantity of matter) of any 
 system is not altered by any process which may take place within that 
 system." Ostwald. This is often designated as the "law of the 
 conservation of matter." (It is frequently desirable to differentiate the 
 changes occurring within a body or a "set of materials" or a group of 
 bodies from those changes which may take place in its "surroundings" 
 To avoid this cumbrous expression, science has introduced the word 
 " system" to designate this assemblage or arrangement of bodies considered 
 as being insulated from its environment.) 
 
 Definite portions of matter are called "bodies" or "specimens." 
 Different kinds of matter are called "substances." Bodies may be 
 homogeneous or heterogeneous accordingly as they are made up of visibly * 
 like or unlike parts. 
 
 Matter appeals to the senses because of its attributes or properties 
 which are characteristic of a body or a substance. 
 
 By the word "property," the idea of a peculiar quality of a thing 
 is conveyed;- - I> .implies that, which is essential and inherent in a thing, 
 as sweetness' Is a' property; of :sugar. "The properties of an object are 
 all the relations in, whjciu it ,can be ( made to appeal to the senses." 
 Ostwald.//- */ ;[ { r ; ' J f r ,\ f'-, , r ; 
 
 The identity of a body or a substance depends upon a definite assem- 
 blage of properties. 
 
 The matter of the universe is continually undergoing a change, i.e., 
 the properties are being altered in degree or completely changed. Any 
 directly observed change taking place in matter is called a phenomenon. 
 
 If the change alters the properties of a substance but temporarily, it is 
 said to be a physical phenomenon; if the properties are changed per- 
 manently, i.e., the substance loses its original identity, the observed 
 change is called a chemical phenomenon. 
 
 Those properties which are exhibited by physical phenomena are 
 called physical properties, and those which require chemical phenomena 
 for their exhibition are called chemical properties. 
 
 (Physics has sometimes been defined as the study of physical phe- 
 nomena, and Chemistry as the study of chemical phenomena.) 
 
 Physical properties as a class are further subdivided into two groups, 
 general and specific properties. General properties are those which 
 are possessed in common by all kinds and conditions of matter. Specific 
 properties are those which are characteristic of a particular kind of 
 matter (a substance), yet common to a given species of matter. Note 
 that "bodies" cannot always be differentiated from one another by the 
 sole use of specific properties, for if they be "bodies" or "specimens" 
 of one kind of matter, say sulphur, then all the bodies will possess in 
 common the specific properties. Observe that such attributes as size 
 
 * Colloidal solutions represent an exception to this rule if the word " visibly " is 
 used in the ordinary sense of the term. 
 
FUNDAMENTAL CONCEPTS. 1 9 
 
 and form assist in differentiating "bodies," yet said attributes do not 
 enter into our concept of the substance of which the body is composed. 
 
 Experiment I. Matter, Bodies and Substances. 
 
 (Record all data in your notebook.) 
 
 Place two small pieces of each of the following upon the top of your 
 laboratory desk zinc, iron, sulphur and glass. How many "bodies" 
 are represented? How many substances? How did you differentiate 
 the bodies? The substances? Name the properties which you made 
 use of in each case. 
 
 A piece of sulphur (one cubic cm.), under the ordinary conditions of 
 the laboratory possessed the following physical properties: color, yellow; 
 odor, practically none; hardness, 2.5; melting point, 115 C.; boiling 
 point, 448 C.; specific gravity, 2; specific volume, .5; specific heat, 1.8; 
 insoluble in water; soluble in carbon bisulphide; crystalline structure; 
 opaque; poor conductor; weight, 2 grams; form, cubical; heat capacity, 
 1.8 x 2; inertia; indestructibility; extension; porosity; impenetrability. 
 
 Do all of the properties enumerated above belong to all bodies? 
 To all substances ? Classify all of the above properties under the three 
 heads "general," "specific," and "body" properties. 
 
 Which class, or classes of properties enter into your concept of matter? 
 Bodies? Substances? 
 
 Can you differentiate bodies by the use of such attributes as weight, 
 form and size? Try to do so by placing the two pieces of sulphur side 
 by side. Can you conceive of any exception to your conclusion ? If so, 
 explain. 
 
 PHYSICAL PROPERTIES. 
 Experiment II. Determination of Specific Gravity. (Quant.) 
 
 (a) Clean and thoroughly dry a 50 cm. 3 Erlenmeyer flask; weigh it 
 accurately and record weight of same in your notebook; from a burette 
 containing distilled water at the temperature of the laboratory, measure 
 into flask 10 cm. 3 of the water. (Recall former instructions relative to 
 meniscus). Weigh as quickly as possible. Calculate the weight of 
 i cm. 3 of water at the temperature of the laboratory. The weight of i cm. 3 
 of a substance is called its density. The weight of the volume of a 
 substance compared with the weight of an equal volume of water at the 
 temperature of its maximum density (what is this temperature ?) is called 
 its specific gravity, i.e., the ratio is given this name. The reciprocal 
 of density or specific gravity of a body, or the volume occupied by one 
 gram, is called its specific volume. 
 
 (b) (Quant.) To determine the specific gravity of glass suspend 
 a piece of silk thread from the hook on the balance and adjust balance 
 so that pointer makes vibrations of equal length on either side of the zero 
 point; tie the thread to the piece of glass so that when a beaker of glass 
 is placed under it, the solid will be completely immersed in the water; 
 
20 EXPERIMENTAL CHEMISTRY. 
 
 weigh the suspended glass, first in the air and then when immersed. 
 Record the two weights in your notebook. What is the difference in 
 weight ? What is the weight of the volume of water equal to the volume of 
 the glass? What is the specific gravity of glass? Its density? The 
 volume of the piece of glass? Its specific volume? 
 
 (c) (Quant.) Determine the specific gravity of alcohol by means of 
 a specific gravity bottle or an Ostwald-Sprengel pyknometer. (Instruc- 
 tions from assistant.) 
 
 (d) (Quant.) To ascertain the specific gravity of a solid when only 
 small pieces are available weigh the specific gravity bottle (Richard's 
 form preferred) when empty; when filled with water and properly ar- 
 ranged; weigh the solid; place the solid. in the bottle and add water until 
 bottle is filled, observing all precautions; weigh. Weight of contents? 
 Of solid in bottle? Of water now in bottle? Does the solid displace 
 its own volume of water? What is the weight of this volume of displaced 
 water? Using small pieces of glass, determine its specific gravity. 
 
 Experiment III. Compressibility of Gases. Boyle's Law. (Quant.) 
 
 To determine the relation between volume and pressure. Plot graph 
 of results. (Instructions from assistant.) 
 
 Experiment IV. Expansion of Gases With Increase of Temperature. 
 Charles' Law. (Quant.) (Instruction.) 
 
 Experiment V. Melting and Boiling Points. (Quant.) (Instructions.) 
 See Traube's "Physico-chemical Methods." 
 
 ENERGY. 
 
 A body may possess other qualities or "conditions," such as motion, 
 electrical charges or temperature. These conditions may be regarded 
 as energy relations of matter. They do not enter into our concept of a 
 substance, yet their influence upon its properties is very marked in some 
 cases. 
 
 Energy is the capacity for work. " It is the essential thing in the 
 universe." Richards, T. W. By some it has been called the funda- 
 mental property of the conceived entity, called "matter." This is prob- 
 ably due to the fact, as science holds, that matter is always associated 
 with more or less energy. The idea of considering energy as a property, 
 thereby subordinating it to matter, is severely criticised by many, who 
 contend that it should be placed on the same plane with matter. The 
 changes in properties and the power to produce them are, therefore, con- 
 ceived to arise not from a number of distinct entities, but from a single 
 one, which is capable, however, of manifesting itself in a variety of differ- 
 ent ways. "That which gives rise to the changes in the properties of 
 bodies and to the power to produce such changes, is called energy." 
 Noyes, A. A. 
 
 The familiar forms of energy are, heat energy, light energy, electrical 
 energy, gravitational energy, kinetic energy, etc. Careful experimentation 
 covering a long period of time has shown that equal amounts of the various 
 
FUNDAMENTAL CONCEPTS. 21 
 
 forms of energy will produce, on transformation, equal amounts of heat. 
 This has led to the enunciation of the law called "conservation oj energy" 
 which may be stated as follows: 
 
 In a limited system, regardless oj the transformations which may take 
 place within the system, the total amount oj energy is not altered. 
 
 Julius Robert Mayer, in 1842, was the first to formulate the above 
 law. It is sometimes called the First Law of Energetics. 
 
 The importance of the role of energy in Chemistry gradually becomes 
 more clear as the individual investigates. It is generally admitted that 
 "the existence of matter is inferred only from its various energy mani- 
 festations." This leads to the conclusion, that at least so far as our ex- 
 perience is concerned, matter and energy are inseparably connected; 
 that a " body" as we observe it, represents so much of the entity, " matter," 
 associated with more or less of the entity, " energy." Matter * never exists 
 alone, i.e., isolated from energy, and science is quite sure that if the energy 
 were removed from a piece of iron, that its properties would be very 
 different. A body or an object possesses in addition to matter a certain 
 " energy content," which determines its properties. This form of energy, 
 we call internal energy. (This does not include kinetic energy.) 
 
 Experiment I. Forms of Energy. Transformation of Energy. 
 
 (a) Raise the object called a "mortar" from your desk. Did 
 this require energy? Work? What kind of energy did you exert? 
 What kind of energy did you work against ? What became of the 
 energy you expended? Was there any energy destroyed in the above 
 operation ? Consider yourself and the mortar a body with its capacity 
 for gravitational energy, as constituting a "system;" did the system gain 
 or lose any energy? Was there a transformation of energy during the 
 process? State specifically the transformation. 
 
 (b) Procure a hard rubber or glass rod and a piece of flannel from the 
 assistant; observe that it is somewhat "cold" to the touch and that if 
 touched to pieces of thin paper it does not disturb them; now vigorously 
 rub the rod with the flannel; hold the rod close to the pieces of paper; 
 what is the effect? Touch the rod and note whether it seems to have 
 changed in temperature. The "rubbing" required the expenditure of 
 what kind of energy ? It was transformed into what two forms of energy ? 
 What bodies constitute our system in this case ? Did the system gain or 
 lose energy? Explain. Are chemical or physical phenomena involved 
 in this experiment? Give reason for your answer. 
 
 (c) Ask the assistant for a d'Arsonval galvanometer, and pieces of 
 iron and copper wires about 20 cm. long. Join the iron and copper wires 
 at one end by twisting them firmly; attach the free ends of the wires to the 
 two binding posts of the galvanometer; take the reading. (Instructions.) 
 Light a match and apply the flame to the juncture (twist) of the two 
 wires; take the reading while applying the heat; this latter form of energy 
 
 * There is much data at hand which points to the conclusion that "matter" is 
 nothing but energy. 
 
22 
 
 EXPERIMENTAL CHEMISTRY. 
 
 is partially transformed into what two forms of energy revealed by the 
 conduct of the galvanometer? 
 
 (d) (Instructions.) Into a cold dilute solution of hydrochloric acid dip 
 a strip of zinc and a strip of platinum; by means of wires connect strips 
 with the binding posts of an ammeter. Result ? Place your hand upon 
 the vessel containing the acid and metals; any change in temperature? 
 Is the zinc being consumed, i.e., is it dissolving in the acid? Has that 
 portion which has been consumed lost its identity, i.e., do you see anything 
 in the acid which seems to possess the same assemblage of properties that 
 zinc possesses ? Those phenomena in which there is involved a change 
 of identity of a substance are called chemical phenomena. As the result 
 of these chemical phenomena, has energy become 
 available for transformation into heat energy and 
 electrical energy? Enumerate as many different 
 forms of energy as you can. 
 
 From a study of the foregoing experiments the 
 student should readily deduce that energy manifests 
 itself in various forms, many of which are familiar. 
 A form of energy with which the chemist is almost 
 continuously concerned is the one illustrated in the 
 last experiment the change of identity suffered by 
 zinc when it dissolved in the hydrochloric acid. The 
 zinc strip is supposed to contain a given amount of 
 and under definite conditions, a 
 further asserted that this internal 
 
 FIG. i. 
 
 zinc matter 
 
 definite quantity of energy. It is 
 energy which is " stored up " in the zinc substance is liberated, at least 
 in part, as "free energy" when the zinc presses into solution, and that it 
 is this so called jree energy which is transformed into heat and electricity. 
 This relationship of matter and energy will receive attention as we ad- 
 vance in our work. 
 
 Experiment II. Relation of Energy to the Properties of Matter. 
 
 (Postulating that our senses can respond to only energy manifestations, 
 proceed with the following experiments.) 
 
 (a) Procure a cylinder of sulphur (or hard rubber or glass rod) and a 
 piece of flannel from the assistant; examine the piece of sulphur; why 
 does it appeal to the senses? Do these properties bear any relation to 
 energy? (Answer in the light of the above postulate.) If they do not, 
 should properties alter with change of " energy content ?" Test your 
 conclusion by experiment; hold the sulphur rod near small pieces of 
 thin paper; has it the property of attracting the paper? Rub the rod 
 vigorously with the flannel; now hold rod near the paper. Is there 
 any alteration or change in the properties of the rod ? Did you expend 
 energy in rubbing? If so, what became of it? Explain in terms of 
 energy relations. 
 
 Rub the rod again with flannel, supposedly "storing up" an electrical 
 
FUNDAMENTAL CONCEPTS. 23 
 
 charge on the sulphur; is its energy content thereby changed? Has 
 this changed its identity? Is it possible that many of its properties are 
 altered in so small degree that we are unable to detect the change ? Does 
 an alteration in small degree of the properties occasion a loss of identity ? 
 
 If a small change in the energy condition produces a small change in 
 the properties of a body or a substance as in the case of sulphur, might we 
 be led to assume that great energy changes will be followed by correspond- 
 ingly great alterations in properties of bodies, perhaps to the extent of 
 changing the identity of a substance, i.e., nearly a complete alteration of 
 properties ? 
 
 (b) Examine the filament in the incandescent lamp over your desk; 
 note its properties as best you can by looking through the glass; observe 
 especially its color, its diameter and its light-giving properties. In view 
 of your previous experience, would you say that the filament possesses 
 a definite amount of internal energy ? Does the filament possess a defi- 
 nite assemblage of properties under the present conditions of temperature ? 
 
 Turn the switch permitting the electricity (energy) to flow through the 
 filament. This energy is transformed into what two forms of energy 
 easily observed? Does the filament contain more or less energy now 
 than when current is not traversing it ? Are the properties of the fila- 
 ment altered when the current is flowing through it? Does it lose 
 its identity either temporarily or permanently? Cut off the current, re- 
 storing the filament to the original " energy content," i.e., to the same 
 conditions of energy. Does it regain its original properties ? If so, what 
 kind of a phenomenon have you been observing? 
 
 As the result of this experiment and those which have preceded it, 
 would you feel justified in affirming that there is a causal relation between 
 energy and the properties of matter? Should the energy conditions be 
 stated before an attempt is made to describe the properties of a substance ? 
 Why? 
 
 It seems quite certain that there is a very intimate relation between the 
 properties of a substance and the amount of energy associated with it, 
 that the properties vary in degree as the amount of internal energy 
 varies. It is quite likely, if we could remove the greater portion of the 
 internal energy possessed by sulphur without resorting to chemical proc- 
 esses, that we should have a substance whose properties would be very 
 different from those of sulphur. 
 
 The above experiment with the filament of the incandescent lamp 
 emphasizes the fact that when we speak of a substance possessing a par- 
 ticular identity, i.e., a definite set of properties, we mean under certain de- 
 fined conditions of energy. 
 
 Experiment III. Relation of Energy to Phase of Substance. 
 
 Examine several small pieces of ice, it is said to be water in the solid 
 state or phase. Has it a definite identity ? Place a few small pieces of the 
 ice in a test tube and gently warm the tube and its contents in a flame 
 
24 EXPERIMENTAL CHEMISTRY. 
 
 (heat energy). As long as ice and water are together in the test tube the 
 temperature of the mixture does not alter. What becomes of the energy 
 that is being contributed by the flame ? When all of the ice has melted 
 remove the test tube from the flame; you now have water in the liquid 
 phase. Does this same weight of water contain more or less internal 
 energy than the equal weight of ice ? Have the properties been altered ? 
 If you had never seen water in the liquid phase, would you have recognized 
 it by virtue of your knowledge of the properties of ice ? Has the identity of 
 the ice been lost ? Introduce the tube and its contents into the flame again, 
 and heat until all of the latter has passed into the gaseous phase. Can 
 you see steam ? Does the steam possess more or less internal energy than 
 the water? Are its properties different from those of water, i.e., has the 
 steam an identity of its own? Are these phenomena of a chemical or 
 physical nature? Would you say that the energy content determines, 
 at least in a large measure, the phase or state of aggregation of water? 
 Which phase of a substance has the maximum amount of internal energy? 
 
 A solid is frequently said to possess "form energy," by which expres- 
 sion we understand that sufficient attraction (cohesive energy) exists 
 among its particles to give it rigidity and form. Because of these attri- 
 butes due to its energy "condition," a solid is sometimes denned as pos- 
 sessing the properties of specific volume and form. A liquid, possessing 
 less effective jorm energy, has not the property of a definite form, but 
 that of specific volume only. A gas has neither the specific property of 
 form nor definite volume; it possesses relatively the minimum of effective 
 form energy. 
 
 When the ice in the above experiment was heated, sufficient energy was 
 imparted to it to overcome the effects of the form energy, at least in part, 
 with the result that the solid was changed into the liquid " state of aggre- 
 gation." "This latter expression came into use because it is commonly 
 assumed that in the different states, the particles of which a substance is 
 composed, are differently aggregated." Heat energy, when imparted to 
 these particles, acts like a " repellent force" overcoming the form energy 
 and causing a change of state of the substance. A similar course of 
 reasoning is followed in explaining a change from the liquid to the gaseous 
 state. 
 
 If form energy is dominant in a substance, then the solid state results; 
 if heat energy is dominant, then the gaseous state prevails. 
 
 The energy content of a substance under definite pressure determines 
 its physical state or phase. Considerable confusion has grown out of a 
 careless use of these last two words. 
 
 PHASES AND STATES OF MATTER. 
 
 If we consider ice, water and water vapor (gas) as constituting a system, 
 then the three physically distinct parts of this system represent three 
 phases, also three states solid, liquid and gaseous. What is the dis- 
 tinction between "phase" and "state?" Ordinarily we consider but 
 
FUNDAMENTAL CONCEPTS. 25 
 
 three states of matter, namely, solid, liquid and gaseous, no difference 
 whether we are speaking of and comparing various substances or just one. 
 The word phase is usually defined as a homogeneous aggregation or mass. 
 Homogeneous means "like," i.e., possessing same properties; heterogene- 
 ous expresses the idea of "unlike." If kerosene were poured upon a 
 little water in a test tube and two distinct layers were seen, how many 
 phases would be present ? How many states ? It is readily seen that we 
 have a mass of -water and kerosene which is in the liquid "state," but 
 that we have two phases, i.e., two masses which are each homogeneous. 
 Aggregations of matter which are visibly different or can be mechanically 
 separated from one another are called phases. Examine a piece of granite 
 and there will be visible three kinds of matter physically different, and 
 which can be separated by a mechanical process, hence we have three 
 phases and but one state (solid), represented. The granite is not a homo- 
 geneous substance; ice is. The student should remember that the num- 
 ber of different substances present does not necessarily determine the 
 number of phases. When you make a solution of salt and water you have 
 but one phase, because the entire mass is homogeneous. Consider a vol- 
 ume of air; it is composed of many substances, nitrogen and oxygen, chiefly, 
 yet there is but one phase, the gaseous. 
 
 Homogeneity must always be the test. The terms phase, homogene- 
 ous and heterogeneous are customarily restricted to usage in the physical 
 sense. 
 
 Experiment IV. A Study of Phases of Matter. 
 
 Place 5 cm. 3 of water in a test tube; how many "states?" Phases? 
 In answering this question ignore air and watery vapor above surface of 
 water. Add 3 cm. 3 of alcohol; shake tube well. You have a mass of 
 alcohol and water; is it homogeneous? How many "states" present? 
 Phases ? 
 
 Repeat above, using kerosene instead of alcohol. Answ r er above ques- 
 tions in order. 
 
 Place 5 cm. 3 of water in a test tube; add 2 grams of salt; shake well and 
 heat gently for a few seconds. How many phases? States? Add 5 
 grams of salt to the solution; treat as before. Do you find that all of the 
 salt will not go into solution? How many states? Ho\v many phases? 
 
 From the above it is evident that the test of " homogeneity" (used in 
 the physical sense), applied to a mass solid or liquid will usually answer 
 the question as to the number of phases present. The physically dis- 
 tinct parts of a system represent the phases of it. There is but one phase 
 recognized in the gaseous "state," but many in the solid and liquid states. 
 
CHAPTER II. 
 CHARACTERISTICS OF PHYSICAL AND CHEMICAL CHANGES. 
 
 Those phenomena in which matter undergoes a temporary change of 
 identity, regaining its original identity when the original conditions of 
 energy are again established, are called physical changes. When the 
 change in identity is permanent, it is known as a chemical change. As a 
 rule, chemical changes are accompanied by physical changes. Chemical 
 changes are defined as those changes which affect the identity, and 
 produce alterations in the substance under defined conditions of energy. 
 Chemistry is primarily the science of these chemical changes and their 
 attendant phenomena. 
 
 Experiment I. Physical and Chemical Changes. 
 
 (a) Examine the properties of a piece of platinum wire which has 
 been sealed into the end of a glass rod; note its physical properties so 
 carefully that you become sure of its identity. Has it a particular 
 group of properties under the prevailing conditions of energy? Now 
 hold the platinum wire in the Bunsen flame for a few seconds. Note 
 its properties w r hile in the flame. Are they altered? Remove wire 
 from flame, does it continue to glow ? Does it emit light now ? When 
 you are confident that wire is under the original energy conditions, ex- 
 amine its properties. Has it lost its original identity ? What kind of a 
 change did the platinum wire undergo ? Define. 
 
 (b) Note the properties of a piece of magnesium " ribbon;" by means 
 of a pair of pinchers, hold the ribbon in the flame for a few seconds and 
 then remove it; does it give out light and heat after it is removed from 
 the flame ? Examine what is left of the ribbon, i.e., the white powder, 
 when it is cooled to the original external energy conditions; has its 
 identity been changed? What kind of a change did the magnesium 
 ribbon undergo? Define. 
 
 Magnesium + Oxygen * Magnesium Oxide. 
 
 In the above experiment it must be evident to the student that more 
 energy is yielded by the burning of the magnesium ribbon than was 
 supplied to it by the flame. Is it possible that when it combined with 
 oxygen, they both gave up a quantity of energy which was transformed 
 into heat and light ? If so, would the energy content of magnesium 
 oxide be less than the sum of the respective energy contents of magnesium 
 and oxygen? 
 
 What kind of changes are illustrated by the rusting of iron, a falling 
 ball, burning coal, melting of ice, and the souring of milk ? 
 
 26 
 
CHARACTERISTICS OF PHYSICAL AND CHEMICAL CHANGES. 27 
 
 Experiment II. (L. T.) Electrolysis of Water. Change of Identity. 
 
 Make a mixture of acid and water in the ratio of i to 20. (Pour the acid 
 into the water slowly.) Fill a Hoffman apparatus (Figs. 2 and 3) for the 
 electrolysis of water with the above mixture. To each of the stop cocks 
 attach a glass delivery tube by means of pieces of rubber tubing, so that 
 they will dip into a vessel of water, a pneumatic trough preferably. Keep 
 stop cocks closed until at least 15 cm. 3 of gas has collected in the tube 
 which contains the lesser volume of gas. What is the volume in the 
 other tube (approximately) ? Now collect in separate test tubes these 
 
 FIG. 2. Voltameter. 
 
 FIG. 3. 
 
 respective volumes of gases, or portions of them, by "water displace- 
 ment," i.e., fill a test tube with water and invert it with mouth under 
 water. Open stop cock slowly and gas will be forced through delivery 
 tube, displacing water in test tube. Place your thumb over the test tube 
 mouth under water and bring its mouth upward to a burning match. 
 What happened? Repeat this operation with the other tube of gas. 
 Using a glowing splinter repeat both operations. 
 
 The larger volume of gas was hydrogen; the other gas was oxygen 
 In which gas did the splinter burn most vigorously ? 
 
 After thus examining the properties of these two gases, would you say 
 that the water has lost its identity? 
 
28 EXPERIMENTAL CHEMISTRY. 
 
 Water > Hydrogen + Oxygen. 
 
 Water is obviously a more complex substance than either hydrogen 
 or oxygen. The water has undergone a chemical change known as 
 decomposition or analysis. 
 
 Experiment III. Effect of Heat on Mercuric Oxide. 
 
 By means of a folded piece of paper introduce a little mercuric oxide 
 
 (red precipitate) into a dry test tube; determine combined weight of 
 tube and powder by means of chemical balance; 
 note properties of the red powder; heat tube 
 gently. Does color of powder change ? Hold a 
 glowing splinter down in the mouth of tube 
 while heating. Results? Do you recall having 
 worked with a gas which revealed similar proper- 
 ties ? Its name ? Examine the sides of the tube; 
 do they show a metallic luster? What is it ? In 
 view of above phenomena what kind of a change 
 would you say you have observed ? 
 
 FIG. 4. (Smith andKeller). As the powder in the tube cools does it tend 
 to regain its original color? Is this action the 
 
 reverse of the one which takes place when tube is being heated ? Weigh 
 
 tube and contents. Result ? 
 
 Mercuric oxide T* Mercury + Oxygen. 
 
 Of the three substances indicated in the above equation which is the 
 most complex ? 
 
 The mercuric oxide has undergone what particular kind of chemical 
 change ? 
 
 Definite conditions of energy seem to be necessary to secure the de- 
 composition of mercuric oxide. Is it possible that the energy con- 
 tributed by the flame is stored up in the mercury and the oxygen ? 
 
 Experiment IV. Effect of Heating Copper in Air. 
 
 Clean and dry the small porcelain crucible which you will find in the 
 drawer; introduce into it about i gram of powdered copper; by means 
 of chemical balance, find total weight of copper and crucible; place 
 crucible upon a pipe-stem triangle (Fig. 4) and heat with Bunsen flame 
 for 15 to 20 minutes. Is there any change in color? While the crucible 
 and contents are cooling examine the material upon the side shelves, 
 labeled ''copper oxide." Does it bear any likeness to substance in 
 crucible? Place crucible and contents upon balance. Have they 
 gained or lost weight? Explain. Which is the more complex, copper, 
 or copper oxide ? 
 
 Copper + Oxygen * Copper Oxide. 
 
CHARACTERISTICS OF PHYSICAL AND CHEMICAL CHANGES. 29 
 
 Chemical changes like the one above involving the addition of sub- 
 stances to one another are called combinations or syntheses. In the 
 experiment, you arranged conditions favorable for the synthesis of 
 copper oxide. 
 
 Experiment V. Solution Facilitates Chemical Action. 
 
 (a) Place about .5 grm. of tartaric acid in a mortar and pulverize it; 
 add to contents of mortar .5 grm. of sodium hydrogen carbonate (bicar- 
 bonate of soda); note that no chemical action takes place; pulverize the 
 mixture; any chemical action? Transfer a little of the powder to a test 
 tube and add water; does a chemical change take place? 
 
 (b) Add 5 cm. 3 of concentrated hydrochloric acid to a test tube; add 
 the same quantity of ammonium hydroxide to another tube; hold the 
 mouths of the test tubes near to each other. What evidence of a chemical 
 change? Explain. (Both of the above reagents are gases dissolved 
 in water.) 
 
 Chemical action occurring, in which one or more substances are solids, 
 must be largely a surface phenomenon, i.e., occurring on the surface. The 
 extent of surface will determine largely the rate at which the action 
 proceeds. Substances, finely divided expose more surface per unit 
 weight, hence give more intimate contact with one another. Solutions 
 of substances provide a convenient method for securing intimate contact 
 of substances. Gases provide an ideal condition but are bulky and 
 inconvenient to manipulate. 
 
 An interesting application of this principle of determining the rate 
 at which the chemical action shall proceed by determining the area 
 of surface exposed, is utilized in the manufacture of explosives. The 
 explosives are cast into cylindrical sticks and the surface is increased by 
 providing a number of longitudinal holes of varying diameters. 
 
 Experiment VI. Relation of Energy to Chemical Change. 
 
 From the reagent bottle containing silver nitrate, procure i or 2 
 cm. 3 of the solution in a clean test tube. Dip the end of the glass 
 stirring rod into the liquid and, using it like ink, write your name upon a 
 piece of white paper; expose the writing upon paper to direct sunlight 
 for half an hour if need be leave it until next laboratory period. Re- 
 sults? Considering your method of procedure, what factor produced 
 the chemical change? 
 
 EXERCISES. 
 
 (Record all answers in notebook.) 
 
 What form of energy caused the change in the platinum wire ? De- 
 composition of water? Analysis of mercuric oxide? Assisted in the 
 synthesis of copper oxide? Change of silver nitrate? Does energy 
 seem to be as closely identified with chemical changes as it was with 
 physical changes ? Has each one of the above changes involved time or 
 
30 EXPERIMENTAL CHEMISTRY. 
 
 have they taken place " instantaneously ? " (The word " instantaneously " 
 is seldom found in the vocabulary of a scientist what do we mean by 
 this affirmation?) 
 
 What is the chief distinction from the standpoint of energy consider- 
 ations between the conduct of the piece of the magnesium ribbon under- 
 going chemical change and the other chemical changes enumerated 
 above ? Was energy necessary to induce the change ? 
 
 Name the different kinds of chemical changes which you have studied 
 to date. 
 
CHAPTER III. 
 ELEMENTARY SUBSTANCES, MIXTURES AND COMPOUNDS. 
 
 It appears from the foregoing experiments that under suitable con- 
 ditions of energy, matter can be made to combine with other kinds of 
 matter with little difficulty, producing more complex substances, but it 
 has not been possible to continue simplifying matter indefinitely. It 
 has been impossible to go beyond resolving all kinds of matter into 
 anything more simple than about 80 so-called elementary substances. 
 These simple substances have resisted every effort to date to decom- 
 pose them into anything more simple or elementary. Examples of 
 these elementary substances are, platinum, magnesium, mercury, oxygen, 
 copper, gold, silver, iron, hydrogen. Water is not an elementary sub- 
 stance as you were able to decompose it into two simple substances, 
 namely, hydrogen and oxygen. Substances which can be resolved into 
 something more elementary are not elementary substances. It has been 
 suggested (Richards, T. W.) that this limit to convertibility might be 
 called the "Second Law of Matter." 
 
 Experiment I. Examination of the Elementary Substances. 
 
 (L. T.) Ask the assistant to place specimens of the 80 elementary 
 substances on the table, then carefully note the physical properties of 
 the various substances. How many are in the liquid state ? Gaseous 
 state? Solid state? Notice flourine, chlorine, bromine, iodine; is there 
 apparently any gradation in properties (physical) ? 
 
 Observe those which have a metallic lustre. The majority of these 
 are called metals. Record the names of those which do not have this 
 property they are called non-metals. 
 
 Experiment II. Energy Content Affects Properties of Elementary Sub- 
 stances. 
 
 (L. T.) Compare the properties of yellow phosphorous with those 
 of red phosphorous. Perform such experiments as the assistant suggests. 
 Note: Don't handle the phosphorus with the hands or take it into the 
 vicinity oj much heat; cut it under water. 
 
 The above is a chemical change due to change in energy content, 
 causing a new internal rearrangement. Consult lecture notes. 
 
 Experiment III. Mixtures Heterogeneous and Homogeneous. 
 
 (a) Try to dissolve a small quantity of flowers of sulphur in water. 
 After shaking vigorously how many phases have you ? Can you separate 
 
32 EXRERIMENTAL CHEMISTRY. 
 
 the t\vo substances by a mechanical process like filtration (instructions) 
 or evaporation ? Did each substance retain its original properties ? 
 
 (b) Dissolve a small quantity of flowers of sulphur in carbon disulphide. 
 (This latter substance is very inflammable and quite poisonous; keep it 
 away from free flames.) How many phases ? Pour a little of the solution 
 into a beaker; set it in the open window for 15 to 20 minutes or longer. 
 Results? Does this prove that you had a mixture? Why? What 
 kind of a mixture? 
 
 A Mixture is an aggregate of substances in which each substance 
 retains its characteristic properties. The various substances of a mix- 
 ture are called "components." One of the best tests of a mixture is, can 
 the substances be separated by mechanical processes. 
 
 Experiment IV. Mixtures and Compounds. 
 
 Mix thoroughly in a mortar 5 grams of iron filings and an equal weight 
 of sulphur. Place a portion of the contents of mortar on a sheet of paper 
 and pass a magnet near it. Note that the iron is withdrawn leaving the 
 sulphur. 
 
 Introduce a small portion of the contents of mortar into a test tube and 
 add 5 or 6 cm. 3 of carbon disulphide; shake tube vigorously; does the 
 sulphur dissolve and leave the iron in the tube ? Pour a portion of the 
 liquid upon a watch glass; examine after 15 or 20 minutes. Results? 
 What particular name would you apply to the contents of the mortar ? 
 
 Now take the portion that remains in the mortar, and by means of a 
 folded paper introduce it into a test tube until tube is half filled; heat 
 test tube to redness. When the iron and sulphur become sufficiently 
 hot they combine and the mass glows brightly although tube is taken 
 out of the flame. (Source of energy?) Cool and break the tube. 
 Examine portions of contents with magnet. Are its properties the same 
 as those of either iron or sulphur? What kind of a change has taken 
 place ? The new substance formed is called a compound. Did the 
 action continue spontaneously when the flame was removed ? Was 
 heat and light liberated as the result of the chemical action ? 
 
 Iron + Sulphur Iron Sulphide. 
 
 When elementary substances like iron and sulphur combine to form a 
 new substance \vhose properties are characteristic, and in w r hich the 
 properties of the combining substances are merged and lost, we say a 
 compound has been formed. The various substances uniting are called 
 the ''constituents." 
 
 All matter may, for purposes of convenience, be divided into three 
 classes, viz.: elementary substances, mixtures, and compounds. 
 
 Experiment V. Interaction of Compounds. 
 
 Recall the interaction of hydrochloric acid and ammonium hydroxide. 
 
ELEMENTARY SUBSTANCES, MIXTURES AND COMPOUNDS. 33 
 
 Experiment VI. Metathesis A Species of Chemical Change. 
 
 Place a few pieces of potassium iodide in a mortar; observe its color. 
 Now add a small quantity of lead nitrate; observe its properties. Is 
 there any evidence of chemical activity ? Pulverize the mass. Is there 
 any indication that a chemical change has taken place ? What is it ? 
 
 This is a case of "double decomposition" or " metathesis." As shown 
 by the following equation, it is a reaction in which the "factors" are 
 represented by one or more compounds and there is an exchange of 
 position between two or more substances. 
 
 Potassium Iodide -f Lead Nitrate Lead Iodide -f Potassium 
 
 Nitrate. 
 Copper -f Mercuric Oxide - Mercury -f Copper Oxide. 
 
 Metathesis is defined as a reaction in which a substance is transferred 
 from one compound to another. 
 
 Substances which have a chemical effect on one another are said to 
 react; the chemical change is called a reaction; the individual substances 
 which participate in the reaction are called reagents or "factors;" 
 the new substances formed are called the "products." The properties 
 which enable substances to react on one another with varying degrees 
 of activity are among the more important of what are known as chemical 
 properties. 
 
CHAPTER IV. 
 NOTE ON THE ENERGETICS OF CHEMISTRY. 
 
 THEORY. 
 
 The importance of the role of energy in its relation to matter, and 
 physical and chemical changes has been obvious in all of the preceding 
 work. The student has become familiar with some of the forms of 
 energy which may be associated with matter; viz., light energy, heat 
 energy, electrical energy, magnetic energy and gravitational energy. 
 (In all the discussions, energy of motion of the body as a whole, i.e. 
 kinetic energy, is excluded.) The truthfulness of the First Law of energy 
 has become more apparent. It has been possible to transform energy 
 from one form to another. We know from observation that heat energy 
 supplied by the combustion of fuel, has driven steam engines, which in 
 turn have driven dynamoes supplying electrical energy which has finally 
 been converted into heat energy and light energy. We have understood 
 fairly well this series of energy transformations, but there has been one 
 part which has not been quite so plain, that is, just how and why was 
 heat energy evolved by the combustion of coal in the air? Why was 
 heat energy and light energy developed by the burning of magnesium 
 ribbon in the air, and why should the combining chemically of iron and 
 sulphur manifest similar phenomena ? It was not quite plain why heat 
 energy and electrical energy should make their appearance when strips 
 of zinc and copper were properly arranged in a dilute solution of sul- 
 phuric acid, although it may have been explained by saying that it was 
 due to the chemical reaction of zinc upon the acid. In all of these cases 
 we failed to observe the presence of any of these familiar forms of energy 
 in or upon the various substances before the chemical action occurred, 
 yet we know that energy was evolved, and that it must have come from 
 some form, possibly some other form of energy in the body of system. 
 (First Law of Energy.) 
 
 The question fairly before us is, can we harmonize and interpret these 
 phenomena in the light of our original concept as to the relation of matter 
 and energy? Can we reach some conclusion in regard to the relation 
 which properties, physical and chemical, bear to the " matter content" 
 or " energy content" or both, of various substances? 
 
 As was stated in the beginning of the work, matter, as we know it, is 
 conceived of as consisting of an entity, "matter," always associated 
 with a definite quantity of another entity, known as "energy," under 
 defined conditions. For illustration a piece of the substance iron 
 as we know it, is held to contain so much "iron matter" and so much 
 
 34 
 
NOTE ON THE ENERGETICS OF CHEMISTRY. 35 
 
 "internal energy," i.e., any body under defined conditions has a definite 
 " matter content" and a definite "energy content." We can not wholly 
 separate the "energy" from the "matter." What we really see when 
 observing a piece of any substance, is a particular species of matter, 
 simple or compound, associated with a definite quantity of internal 
 energy. 
 
 That a very intimate relation does exist between the above two enti- 
 ties is evidenced by the experiments which have been performed. An 
 alteration in the energy content of a body or a system produced certain 
 changes in the physical properties of the substances involved. Many 
 times the physical properties seemed to be sort of functions -of 
 the energy content. It was observed also that the energy conditions 
 play a very prominent part in inducing chemical changes. It has been 
 obvious in many of the chemical phenomena, if the Law of Conversation 
 of Energy is universal in its application, that in addition to the change 
 of matter, there has been not only a material altering in the quantity 
 of the internal energy of the bodies or systems undergoing such change, 
 but the energy has been frequently transformed. 
 
 A more definite conception as to what we mean by the word " internal 
 energy" or "energy content" may aid in an interpretation and explana- 
 tion of the mechanism of the phenomena all of which at present are not 
 understood. 
 
 By the expression internal energy, is meant the total energy, regardless 
 of form, within a body or system. (This does not include kinetic energy.) 
 In addition to the forms of energy already enumerated, and which may 
 be associated with matter, there are other forms even more uniformly 
 associated with matter, namely, cohesion energy, disgregation energy 
 and chemical energy. The first two are those forms of energy which a 
 body "possesses in virtue of the tendency of its particles to approach 
 and to recede from one another respectively." In solids the cohesion 
 energy is relatively greater than the disgregation energy; the reverse is 
 true of gases, where the particles tend to repel one another. Elasticity 
 of matter, surface tension, contraction of substances on cooling, etc., 
 all are manifestations of cohesion energy. When a body undergoes 
 change of state by the application of heat energy, that which is directly 
 accomplished is the supplying of sufficient disgregation energy to over- 
 come, at least in part, the cohesion energy. It is evident that the 
 internal energy of the body has been increased, i.e., the energy content 
 is now greater than it was before the change of state. It follows, owing 
 to this relation between cohesion energy and disgregation energy, that 
 if this body after change of state were placed in such a relation that it 
 did not continue to receive heat energy, it would give out heat to its en- 
 vironment and the internal energy would be diminished in amount. Any 
 change in the relation of these two forms of energy in a body will be followed 
 by a change in the internal energy, i.e., heat energy will either be absorbed 
 or evolved. Physical changes, including change of state, receive a 
 comparatively satisfactory explanation by a line of reasoning simila 
 
36 EXPERIMENTAL CHEMISTRY. 
 
 to the foregoing, but not so with the majority of chemical changes. The 
 question might be asked here, what "caused" the magnesium and 
 oxygen of the air to react on one -another? Were cohesion energy and 
 disgregation energy the real "cause" of the reaction? Science says no. 
 Again, was the energy evolved by the combination of magnesium and 
 oxygen exactly equal to the diminution of either cohesion energy or 
 disgregation energy or both, in the system composed of the above two 
 substances ? It is true that these two forms of energy may have been 
 altered in degree, but there seems to be another form of energy asso- 
 ciated with matter which is probably the main source of energy in the 
 above example, and is possibly the "cause" of the reaction of substances. 
 It is called chemical energy, and is possessed by systems in virtue of 
 "this tendency of the substances which comprise the system, to undergo 
 transformations into other substances." The exact nature of this so- 
 called chemical energy has not been determined, but it may be compared 
 with potential energy, for it appears to depend largely if not altogether 
 upon the relative positions of matter in which it resides. An illustration 
 of this is a mixture of hydrogen and oxygen which seems to possess 
 potential energy, because of the proximity of its particles. If a spark is 
 passed through the mixture a chemical change results with the trans- 
 formation of chemical energy into heat energy. An impulse was required 
 to induce the reaction. (It might be said in a popular way, to overcome 
 "chemical inertia.") This necessity of contributing a small quantity 
 of energy, relatively negligible, to initiate the transformation of energy 
 seems to be characteristic of many phenomena, both chemical and 
 physical. We observed this in the combustion of the magnesium ribbon 
 and in securing the combination of iron and sulphur. Afterwards the 
 reactions proceeded spontaneously. A ball lying at the top of an in- 
 clined plane, requires an impulse to start it. 
 
 Chemical energy is often regarded as being composed of two factors, 
 one unknown, and the other, possibly, chemical affinity. It is obvious 
 that chemical energy constitutes but a part of the internal energy; also 
 let it be noted that in a system (other conditions being the same) where 
 there is considerable tendency of substances to react (affinity) as in the 
 case of hydrogen and oxygen, or magnesium and oxygen, there will be 
 represented more chemical energy, hence more internal energy, than 
 in the case of a system where there is little or no tendency of the substances 
 to react. It is easily seen, then, that if the systems are alike in every 
 other respect, that in the former case, the substances will combine with a 
 greater liberation of energy, due in the main to a transformation of the 
 chemical energy. It can not be deduced from this, however, that the 
 diminution in the internal energy of the system, which can be measured, 
 is a measure of the chemical energy or affinity of substances, for it is 
 possible and quite likely that during the chemical change there have 
 been changes in other forms of energy, such as the cohesion and dis- 
 gregation energies. This latter change would also produce an alteration 
 in the heat-capacity of the system which would give it the property to 
 
NOTE ON THE ENERGETICS OF CHEMISTRY. 37 
 
 liberate or absorb heat, in degree beyond that which the system originally 
 possessed, exclusive of its chemical energy. Other factors enter also to 
 prevent ordinary usage of the "diminution of the internal energy" of a 
 system as a measure of the " affinity" of substances in the system, but a 
 discussion of same would be out of place in this "note." 
 
 The amount of energy available for work, liberated by a body or a 
 system, is called the "free energy" (U), while that which is not available 
 to do work is called "bound energy." The total energy (internal energy) 
 of a system is equal to the sum of the "free" and "bound" energies. 
 It should not be inferred that because the internal energy of a system 
 is large, that it necessarily follows that the "free" energy content is 
 large. The chemical energy which may be liberated and transformed 
 into free energy from any one substance will differ in amount as it reacts 
 with different substances. When sixteen grams of oxygen combine 
 with hydrogen to form steam about 57,000 calories of heat are liberated; 
 when the same amount of oxygen combines with solid carbon to produce 
 gaseous carbon-dioxide, about 48,500 calories are evolved. 
 
 A law of much importance in chemistry is, 
 
 That a reaction proceeds in the direction in which there will be the 
 greatest diminution oj free energy. 
 
 This "free energy" is perhaps most frequently liberated as " heat 
 energy." 
 
 Another generalization with which every student of chemistry should 
 be acquainted, and which will enable him to frequently determine the 
 direction of a given reaction when placed under new conditions of energy, 
 is known as Le Chatelier's Theorem.* 
 
 "Any change in the /actors of equilibrium from outside, is followed by 
 a reverse change within the system." 
 
 APPLICATION OF THEORY. 
 
 The questions proposed regarding the explanation of various chem- 
 ical phenomena remain as yet unanswered. What interpretation can we 
 offer from the view point of the foregoing theory ? Let us take the case 
 of the combination of iron and sulphur. These are elementary sub- 
 stances; they are not elements. The word "element" is restricted to 
 mean "matter" alone, as "iron matter," or "sulphur matter." It is 
 an abstraction. We have had no experience with the "elements;" that 
 with which we have experience are the two "elements" associated with 
 their respective and appropriate loads of energy. If the energy could 
 be removed, what would they look like? They each possess a given 
 amount of internal energy (cohesion energy, disgregation energy, chem- 
 ical energy). The two simple substances were ground together and the 
 mixture placed in a test tube. We conceive of this mixture as consti- 
 tuting a system. Its internal energy is equal to the sum of the energy 
 contents of iron and sulphur. There was no evidence of chemical 
 action, nor anything to indicate the presence of chemical energy, yet 
 when a small quantity of heat energy an impulse was imparted to it, 
 * Bancroft, The Phase Rule. 
 
38 EXPERIMENTAL CHEMISTRY. 
 
 the iron and sulphur combined with the evolution of considerable light 
 and heat, to form a compound, known as iron sulphide. The heat and 
 light were the result of the transformation of chemical energy chiefly, 
 with a possbility that a portion was yielded by the transformation of 
 the cohesion and disgregation energies of the respective substances when 
 they passed into the form of the new substances. 
 
 The compound (iron sulphide) contains the two "elements" but their 
 energy contents are very different from what they were originally, and 
 therefore we expect the properties of the compounds to be quite differ- 
 ent. Is it likely that if we were able to return to the "elements" their 
 original and respective loads of energy that we would secure a return of 
 the elementary substances, iron and sulphur? Consider the decom- 
 position of mercuric oxide. It will be recalled that the reaction con- 
 tinued spontaneously. In most cases where the chemical change pro- 
 ceeds as above to a completion, the free internal energy of the system 
 diminishes by being transformed into other forms of energy. We know 
 that in the combustion of carbon (coal) in air that a very large per 
 cent of the heat of combustion is convertible into free energy (external 
 work). 
 
 In concluding this discussion, it might be well to give an example 
 of a chemical change in which there is no apparent alteration in the 
 " matter," the only change being in the energy content. If yellow phos- 
 phorous is heated in a closed tube out of contact with the air at a tem- 
 perature of 250 C., it will be converted into the "allotropic" form 
 known as red phosphorous with an evolution of heat. 
 
 Yellow Phosphorous. Red Phosphorous. 
 
 Yellow color, Red color, 
 
 Poisonous, Non-poisonous, 
 
 Burns at low temperature, Burns only at high temperatures, 
 
 Phosphorescent, Non-phosphorescent, 
 
 Soluble in C S 2> Insoluble in C S 2 , 
 
 Garlic odor, No odor. 
 
 These two substances so different in properties possess in common 
 the "element" phosphorous, but the internal energy of each is different. 
 
 Elementary substances, which by virtue of different energy-contents, 
 possess different properties, are called allotropic. Other examples are, 
 graphite, diamond and amorphous-carbon. 
 
 THERMOCHEMISTRY. 
 
 That branch of chemistry which deals with the thermal (heat) changes 
 accompanying chemical reactions is known as Thermochemistry. The 
 chief aim of Thermochemistry is to determine the relative "affinities" of 
 different substances by the principle of "the development of the great- 
 est amount of heat." This as ordinarily used furnishes only approxi- 
 mate data. 
 
 The principles of thermochemistry are in part summarized by Bloxam 
 
NOTE ON THE ENERGETICS OF CHEMISTRY. 39 
 
 as follows: (i) Every chemical change is accompanied by a thermal 
 change which is a constant quantity. (2) The thermal change occurring 
 during the combination of elements to form a compound is called the 
 heat of formation* If heat is evolved as it usually is, it is said to be an 
 exothermic compound; if heat is absorbed, endothermic. Endothermic 
 compounds (usually explosives) contain more internal energy than the 
 constituents originally possessed, and are unstable. (3) The thermal 
 change occurring during the decomposition of a compound is called the 
 heat of decomposition. (4) The heat of decomposition is identical with, 
 but opposite in sign to the heat of formation of a compound. 
 
 The measurement of the thermal changes is accomplished by means 
 of a calorimeter. A very accurate form of this apparatus, known as 
 the "adiabatic calorimeter" has been used for several years in the 
 laboratory of T. W. Richards, of Harvard University. 
 
 The two chief laws of thermochemistry are: 
 
 The First Law, advanced by Laplace and Lavoisier before the first 
 law of energy had been formulated, states that: 
 
 As much heat is given out in the formation of a substance as is required 
 to separate it into its constituents. 
 
 The Second Law (Law of Hess) 
 
 No matter how many stages there are in a given chemical reaction, the 
 quantity of heat liberated depends upon the initial and final states. 
 
 This is sometimes referred to as the "Constancy of the Heat Sum." 
 Berthelot proposed a Third Law 
 
 "Reactions go in the direction in which there will be the greatest 
 evolution of heat." 
 
 Whereas this has considerable value as a sort of working "rule," it is 
 not to be considered as a law, for it is now known that "free" energy, and 
 not " total " energy, determines the direction of the reaction. 
 
 *NoTE. The " heat of reaction " is equal to the difference of the sum of the heats 
 of formation of the original substances and the sum of the heats of formation of the 
 final substances the heat of formation of the elements being counted zero. 
 
CHAPTER V. 
 
 SCIENCE ITS METHODS OF DEVELOPMENT ; 
 CLASSIFICATION. 
 
 DEVELOPMENT OF SCIENCE. 
 
 Men have been led naturally to investigate the causes of the various 
 phenomena observed in the physical universe. As the result of these 
 investigations we have our several sciences. 
 
 The initial step in the development of any science is the securing of 
 knowledge by observation and experiment. By an experiment is usually 
 meant the process of exhibiting certain phenomena under conditions 
 proposed and controlled by the experimenter. In the conduct of these 
 experiments and in all investigation, we proceed upon the truth of wh at 
 is now known as a general maxim in physical science "the constancy of 
 nature," i.e., "like causes produce like effects" and that, irrespective of 
 time and place. The facts secured by such procedure, isolated, and 
 apparently unrelated, become the data of science. A fact is that which 
 has been demonstrated with such accuracy as to leave no reasonable 
 doubt of its truthfulness. As the investigation of nature proceeded, 
 the multitude of facts became cumbersome and unwieldy. Men began 
 to sift this diverse and apparently unrelated data, and found much that was 
 common to many phenomena. In the stating of these universal and valid 
 relationships they formulated what are now known as laws. After ob- 
 serving the gravitational tendency of many bodies, which observations 
 in themselves constituted facts, the universal relationship was recognized 
 and enunciated in the "Universal Law of Gravitation." A law is a 
 statement of a constant relation common to a number of phenomena; it 
 is the expression of a relation among facts. " When two substances agree 
 in some few properties they also agree with regard to all other properties." 
 Ostwald. This statement of the relationship, as a summary of ob- 
 served facts, is called a law. It makes no endeavor to explain. The 
 student should understand that "it does not govern or state what will 
 happen; it predicates an invariable relation." Nernst. There are no 
 exceptions to laws; if there is an exception to a law it is no law. The 
 word "rule" has been frequently confused with the word "law" in this 
 latter sense. A "rule" is merely a statement of the usual order of events 
 the customary procedure and condition of things, as, the earth turns 
 upon its axis once a day. "Exceptions to rules do not destroy them as 
 rules, in fact we say 'the exception proves the rule.'" Richards, T. W. 
 
 It is interesting to note the two methods by which these relationships 
 
SCIENCE ITS METHODS OF DEVELOPMENT; CLASSIFICATION. 41 
 
 have been discovered. When the individual, by experiment and obser- 
 vation, gathers data concerning phenomena among which he suspects a 
 relationship, and arranges it empirically and then endeavors by experi- 
 ment to verify his suspicions, he is said to resort to the inductive method, 
 sometimes called the empirical method. 
 
 The other method used in an endeavor to discover relationships is 
 known as the deductive or theoretical method, and is largely a mental proc- 
 ess. It consists in starting with merely a conception regarding the re- 
 lations among phenomena, then by pure speculation to make new de- 
 ductions, which must be verified by research. 
 
 The inductive method leads us to definite and usually reliable results. 
 
 The deductive method is more fascinating and enables us to penetrate 
 more deeply into the relationships, but we must be very careful and 
 cautious in the choice of the original concept which many times can not be 
 verified by a direct test. 
 
 Conceptions of this kind which can not be directly verified are called 
 hypotheses. The imaginative element predominates in these concepts 
 which are mere assumptions of "the existence of conditions of which 
 we can have no direct experimental evidence." An hypothesis is an 
 effort to explain the mechanism, the "why" of certain phenomena, and 
 to correlate same, that from this general proposition new conclusions may 
 be drawn. This new body of conclusions constitutes a theory, which 
 should be tested by experimental methods. A prediction is merely a 
 guess and can be tested. 
 
 In this connection our attention is called to the scientific and systematic 
 methods the two being commonly confused, although they are entirely 
 distinct. Science consists in discovering the laws to which phenomena 
 conform and is exemplified by such studies as chemistry, biology, mathe- 
 matics and similar analytical studies. System consists merely in the 
 classification of phenomena as illustrated by grammar, history, etc. 
 Natural history preceded biology, the former merely classified; the latter 
 has become a true science by the introduction of analytical ideas of 
 the relations of heredity, environment, etc. Many systematic studies 
 have now given away to the truly scientific. 
 
 CLASSIFICATION OF SCIENCE. 
 
 In the course of time knowledge became so extensive and diverse that 
 it was evident that the great bulk of related knowledge which might 
 properly be called Natural Science must be divided for convenience into 
 a number of more limited ones. These various divisions of the one large 
 study or science are called the "natural sciences." The distinction 
 betwe.en "living matter" and "non-living matter" is made the basis for 
 a division of these natural sciences into two great groups known as the 
 Biological Sciences and the Physical Sciences. The Biological Sciences 
 deal with matter in the living condition, while the Physical Sciences has 
 to do primarily with matter in the lifeless condition. 
 
42 EXPERIMENTAL CHEMISTRY. 
 
 Matter. 
 
 I. Organic Biological Sciences. 
 A. Biology (general), 
 a i. Botany, 
 a 2. Zoology. 
 
 II. Inorganic Physical Sciences. 
 
 A. Physics and Chemistry. 
 
 a i. Astronomy (Physics of the heavens), 
 a 2. Meteorology, 
 a 3. Geology, etc. 
 
 As will be observed from the above outline it is possible to extend 
 our original classification into what are known as the General or Specula- 
 tive Sciences and the Special or Descriptive Sciences. Physics and 
 Chemistry occupying themselves \vith the study of general properties and 
 "transformations" of bodies regardless of external form, and dealing 
 with the substance only, are called the Speculative or General Sciences. 
 Geology, Astronomy, Meteorology, etc., considering distinct classes of 
 bodies in reference to form, classification and distinguishing characteristics, 
 are known as the Descriptive or Special Sciences. 
 
 Chemistry is a descriptive science in so far as it considers the external 
 properties of chemical substances. 
 
 Again all knowledge may be gathered under the respective headings of 
 the Abstract Sciences and the Concrete Sciences. 
 
 The Abstract Sciences, including logic, mathematics, etc., are not con- 
 cerned primarily with matter, but with abstract conceptions. The 
 Concrete Sciences are concerned with bodies, i.e., "definite aggregation of 
 matter," whether living or lifeless. Zoology, Astronomy, etc., would be 
 included under this head. Physics and Chemistry, having to do with 
 both the abstract and the concrete, are frequently named the " abstract- 
 
CHAPTER VI. 
 QUANTITATIVE RELATIONSHIPS. 
 
 LAWS AND CHEMICAL EQUIVALENTS. 
 
 As a science progresses in accuracy it progresses in the true scientific 
 spirit. We have been led to the conclusion in the foregoing work that 
 definite quantities of any substance under defined conditions possess a 
 definite quantity of internal energy. The question might rise logically, 
 as to whether the "factors" and "products," as referred to matter, in- 
 volved in a chemical change bear any such definite relation to one another, 
 that is, when substances combine to form new compounds, (a) does the 
 matter in the system sustain an appreciable loss or gain, (b) do the re- 
 acting substances combine in any definite ratio by weight or volume, (c) 
 do the constituents of the new substance bear to one another a fixed ratio 
 by weight or volume, (d) does variation in the conditions cause a variation 
 in any possible ratio in which substances may combine with one another? 
 To aid us in understanding the conclusions which have been reached in 
 regard to these questions, as the result of much patient and painstaking 
 labor, let us perform a series of experiments. 
 
 Experiments are of two kinds, qualitative and quantitative. The former 
 are resorted to in order to illustrate a property or group of properties. 
 We consider only what takes place. Quantitative experiments, for ex- 
 ample, accept the fact that gases are compressible, then endeavor to 
 determine the exact amount of compressibility for each increment of 
 pressure. They seek to determine any possible mathematical relation 
 of cause and effect. 
 
 Experiment I. Conservation of Mass. 
 
 (a) (L. T.) Analysis of Water. Counterpoise an electrolysis apparatus 
 upon a scale pan, or better, fit a rubber stopper containing three holes 
 into a 250 cm. 3 flask; half fill the flask with water slightly acidulated 
 with sulphuric acid; insert two electrodes through holes in cork, reserving 
 third hole for a glass tube with rubber tubing and pinch cock attached ; 
 by means of this latter and an air-pump exhaust air above water from 
 flask ; close pinch cock ; invert flask ; mark height of contents ; counterpoise 
 apparatus on balance ; turn on the electric current ; note bubbles of gas 
 arising from electrodes. Caution: Don't continue process of decompo- 
 sition too long. Counterpoise apparatus again ; any gain or loss in weight ? 
 Is the sum of the weights of the "products" equal to the sum of the 
 weights of the "factors?" 
 
 Water Hydrogen + Oxygen. 
 
 43 
 
44 
 
 EXPERIMENTAL CHEMISTRY. 
 
 (b) (L. T.) Synthesis of Water. Perform this operation by means of a 
 Hoffman eudiometer (Figs. 5 and 6). Introduce into apparatus about equal 
 volumes of hydrogen and oxygen ; observe volume of the mixture of gases. 
 
 Note. It is well to use an excess of one or 
 the other of the gases so that when the con- 
 traction of volume of gases occurs, there 
 will be a cushion of gas between top of glass 
 tube and the mercury. What deductions, as 
 regards the relation of products 
 and factors, can you make ? 
 
 Hydrogen + Oxygen Water. 
 
 As the result of much refined 
 work it has been concluded that 
 regardless of chemical change 
 there is "conservation of mass" 
 in a limited system. It is quite 
 possible that Van Helmont 
 (1577-1644) was acquainted 
 with the principle. The law 
 was first enunciated by Lavoi- 
 sier (1789). It is commonly 
 spoken of as the law of the "in- 
 destructibility of matter." 
 
 Experiment II. Combina- 
 FIG. 5. tion in Definite Proportions by 
 
 Weight. 
 
 (a) Synthesis of copper oxide. Clean and dry a small 
 porcelain crucible; w r eigh it; put into it about i gram of 
 powdered copper; secure the exact weight; place crucible on 
 pipe -stem triangle which may be supported by tripod or ring 
 clamp. Heat crucible and contents to redness in a Bunsen 
 flame for fifteen minutes ; remove flame and let crucible cool ; 
 weigh; make record of weight. Repeat operations until 
 weight becomes constant. 
 
 Copper + Oxygen * Copper Oxide. 
 Tabulate your data as follows: 
 
 Weight of crucible and copper o. ooo 
 
 Weight of crucible o . ooo 
 
 FIG 6. 
 
 Weight of copper o . ooo 
 
 Weight of crucible and copper before heating 
 
 Weight of crucible and copper, after heating 
 
 Weight of oxygen, combined with copper 
 
 Eudiometer. 
 
 o.ooo 
 o.ooo 
 
 o.ooo 
 
QUANTITATIVE RELATIONSHIPS. 45 
 
 Calculate how much oxygen would combine with a gram of copper. 
 What is the ratio of the combining weights of copper and oxygen under 
 above conditions ? What is the class average ? 
 
 (b) Optional. Synthesis of magnesium oxide. Prepare crucible and 
 cover as above; w r eigh into it about 300 mg. of magnesium ribbon; use for- 
 ceps and scissors for manipulation of ribbon; make a record of the exact 
 weights as above. Place covered crucible on triangle and cautiously 
 heat it to a dull redness then quickly removing flame, raise cover of 
 crucible and observe whether the magnesium starts to glow, which is 
 desired, if it does, quickly replace cover so that none of the magnesium 
 oxide escapes. This operation of heating to redness, removing flame 
 and lifting cover must be repeated until a constant weight is secured 
 and the contents of crucible have a white or grayish color. Tabulate data 
 and calculate ratio of combining weights as in (a). 
 
 Magnesium + Oxygen * Magnesium Oxide. 
 
 (c) Optional. Synthesis of iron oxide. Pure iron wire dissolved 
 in nitric acid gives a compound which on being heated sufficiently, is 
 converted into iron oxide. Into a thoroughly cleaned and dried evap- 
 orating dish which has been weighed, introduce about 750 mg. of iron 
 wire. (Instructions.) Record exact weight of wire used. Place cover 
 glass on dish; add 10 cm. 3 of dilute nitric acid. Now fill a beaker nearly 
 half full of water and heat the water until it boils vigorously ; place covered 
 evaporating dish on beaker so that the latter serves as a steam bath ; keep 
 cover on until all of the iron wire has dissolved, then remove cover and 
 evaporate to dryness; the direct flame may now be applied to the dish; 
 continue its use until a constant weight is secured. As long as red fumes 
 are observed rising from the dish, it will be unnecessary to cool and 
 weigh. The powder remaining in dish is iron oxide. Tabulate data and 
 calculate as in (a). What is the ratio of the combining weights ? 
 
 It is quite obvious in these experiments that we have not been able to 
 get substances to combine in any promiscuous ratio; on the contrary the 
 combining substances have united in certain fixed ratios by weight. A 
 generalization of this principle is known as the "Law of Definite Pro- 
 portions." It was practically confirmed and announced by Proust in 
 .1806, and came largely as the result of a controversy between Berthollet 
 and Proust. The law as frequently stated is, 
 
 The same compound always contains the same elements combined together 
 in the same proportion by weight. 
 
 Experiment III. Constancy of Composition. 
 
 Liberation of oxygen by decomposition of potassium chlorate. 
 
 Note. A small quantity of potassium chlorate may be heated in a test 
 tube and the evolved gas tested for oxygen. 
 
 Clean and dry and heat a porcelain crucible and lid ; cool in a desiccator; 
 weigh accurately and record weight ; introduce into the crucible an accu- 
 
46 EXPERIMENTAL CHEMISTRY. 
 
 rately weighed quantity, say 1.2 gram, of powdered potassium chlorate 
 which has been previously dried at a temperature of about 100 C. for 15 
 to 20 min. Place crucible and contents on a pipe-stem triangle and heat 
 gently in order to prevent spattering ; if any of the salt is deposited upon the 
 cover, the latter must be cooled, and the salt returned to crucible. When 
 the mass solidifies and action has apparently stopped, increase the heat 
 until a perfectly white, non-crystalline mass which is not altered by a 
 further increase of heat, results. Remove cover and heat strongly for a 
 few minutes; cool in desiccator and weigh as before. Record weight. 
 Heat again for 10-15 min.; cool; weigh. Repeat until weight is practi- 
 cally constant. Tabulate all data as in a previous experiment. Calcu- 
 late the percentage of oxygen in potassium chlorate. 
 
 Above result indicates in the case of this compound, potassium chlorate, 
 that a definite proportion by weight of oxygen can be secured from a defi- 
 nite weight of the compound. Further experimentation confirms this 
 principle, known as the "law of constancy of composition." It is but 
 a special statement of what law ? 
 
 Experiment IV. Combination of Two Substances in Varying Propor- 
 tions by Weight. 
 
 (a) Optional. Synthesis of mercuric iodide (Hg I 2 ), as compared with 
 synthesis of mercurous iodide (Hg I). (Instructions from the assistant.) 
 
 (b) A qualitative experiment designed to show in a general way the 
 difference between the properties of two compound substances, each of 
 which is composed of the same elementary substances but in different 
 proportions by weight. Synthesis of nitrogen dioxide (N 2 O 2 ) as com- 
 pared with the synthesis of nitrogen tetroxide (N 2 O 4 ). Assemble parts 
 of a gas generating flask equipped with delivery tube ; place about 10 grams 
 of copper turnings in the flask; arrange delivery tube so that it dips into 
 a vessel of water (pneumatic trough). Pour through the funnel tube 
 about 15 cm. 3 of nitric acid diluted with its own volume of water, then 
 add a few cm. 3 of concentrated nitric acid at a time until there is a rapid 
 evolution of gas; after gas has been evolved for a few minutes collect 
 several bottles of the colorless gas by water displacement; place a glass 
 plate or a piece of wet paper over mouth of bottle on lifting it from water; 
 remove cover so that oxygen of air can come into intimate contact with 
 the colorless gas. Does the gas in the bottle become brownish-red in 
 color ? The colorless gas is a combination of nitrogen and oxygen in the 
 ratio by weight of 14 to 16; the brownish-red gas is composed of the same 
 gases but in the ratio by weight of 14 to 32. Were the conditions under 
 which two compounds formed identical ? Explain. 
 
 This is an illustration of the "Law of Multiple Proportions." 
 // two elements unite in more than one proportion, forming two or more com- 
 pounds, the different weights of one oj the elements, which in the different com- 
 pounds are united with identical amounts oj the other, bear a simple ratio 
 to one another. 
 
QUANTITATIVE RELATIONSHIPS. 
 
 47 
 
 It is not an exception to the law of definite proportions which applies 
 to either one of the compounds formed. The conditions under which 
 each compound was formed were different. See under " Copper." Exp., 
 Preparation of Cuprous and Cupric Chlorides. 
 
 Experiment V. Combining Volumes of Gases. 
 
 (a) Synthesis of water from hydrogen and oxygen gases. Introduce 30 
 cm.3 of each of the two gases into a Hoffman eudiometer, U-form (Fig. 5). 
 Observe proper precautions to collect gases at atmospheric pressure. 
 Open lower stop cock permitting some of the mercury to run out, thus al- 
 lowing gases to expand; pass spark, and then bring mercury columns 
 to same level. Read. Repeat using 25 cm.3; 2 o cm.3, of each gas. R C _ 
 cord all data. What is the average of the ratios of the combining volumes ? 
 
 FIG. 7. 
 
 Note. Before passing spark through mixture of gases it is well to place 
 thumb on open tube to prevent the mercury from being forced out. 
 
 (b) Synthesis of steam from hydrogen and oxygen gases. 
 
 Introduce 25 cm. 8 of each of the two gases into an eudiometer (Fig. 7) 
 which is inclosed in a glass tube which may serve as a steam jacket ; the eudi- 
 ometer should be connected'with a leveling bulb and the gases collected over 
 mercury; the gases must be perfectly dry, as should the mercury; after 
 gases have been collected, pass steam through jacket until entire system 
 is at the temperature of steam ; observe volume of mixture of gases when 
 mercury in eudiometer is at same height as it is in the leveling bulb; 
 lower leveling bulb permitting gases to expand ; clamp tubing below eu- 
 diometer; pass spark through mixture of gases; bring mercury columns 
 to the same height again; note volume. Deduct one-fourth the volume 
 of the mixture of gases at 100 C. from this final volume (Why?). What 
 relation does this volume bear to three-fourths of the volume of the mix- 
 ture of gases at the temperature of the steam ? 
 
48 EXPERIMENTAL CHEMISTRY. 
 
 Note. If more desirable to teacher, the exact combining volumes 
 may be used as prepared by electrolysis of water. 
 
 Experiment has shown that two volumes of hydrogen gas combine 
 with one volume of oxygen gas to produce water. Further it has been 
 shown that when the quantity of water has been converted into steam, 
 that its volume is to the sum of the combining volumes of hydrogen and 
 oxygen at the same temperature, as 2 to 3. 
 
 These and similar experimental results have given rise to a law, known 
 as the "Law of Gaseous Volumes" or "The Law of Gay-Lussac. " 
 
 When chemical action takes place between gases, either elements or com- 
 pounds, the gaseous product bears a simple relation to the volume of the 
 reacting gases. 
 
 A relationship which may be noted at once by a study of the foregoing 
 experiments is stated as the Law of Reciprocal Proportions, or Law of 
 Equivalent Proportions. 
 
 The weights of different elements which combine separately and with one 
 and the same weight of another element, are either the same as or are simple 
 multiples of, the weights oj these different elements which combine with 
 each other. 
 
 It might be called the "law of inter-equivalence of equivalent weights." 
 An application of this principle may be made as follows: 63.6 parts by 
 Aveight of copper or 24.3 parts by weight of magnesium combine with 1 6 parts 
 by weight of oxygen; 2 parts by weight of hydrogen also combine with 1 6 
 parts by weight of oxygen ; therefore, it is obvious that 63.6 parts of copper, 
 24.3 parts of magnesium and 2 parts of hydrogen are chemically equiva- 
 lent, that is, these proportions by weight satisfy the chemical affinity of one 
 another. Again it is to be observed that i part by weight of hydrogen is 
 chemically equivalent to 31.8 parts by weight of copper, 12.2 of magne- 
 sium, and 8 of oxygen. Inasmuch as hydrogen, the lightest of all ele- 
 ments, has the smallest combining weight, Dalton proposed to make 
 this weight the standard, and express all combining weights in terms of it. 
 Berzelius proposed oxygen, which is about 16 times as heavy. These 
 weights of elements which are chemically equivalent to i part by weight 
 of hydrogen are called chemical equivalents. More accurately, that weight 
 of an element which combines with or displaces 8 parts of oxygen or 1.0075 
 parts of hydrogen by weight is called the equivalent weight. Specifically, 
 it is the number of grams oj an element which combines with or replaces 
 1.0075 grams of hydrogen or 8 grams oj oxygen. 
 
 IF the student finds this word "equivalent" difficult to understand in 
 this connection, some satisfaction may be secured by knowing that this 
 same word delayed the development of chemical theory many years. 
 
 Experiment VI. Equivalent Weight of Zinc by Displacing of Hydrogen. 
 
 See Exp. under "Hydrogen." 
 
CHAPTER VII. 
 OUTLINES. 
 
 THEORIES, FORMULAE, VALENCE, EQUATIONS AND UNITS. 
 
 Structure of Matter. j Continuous or discontinuous? Matter bears 
 evidence of possessing a granular structure. 
 
 Statement and Historical Development of the Kinetic-Molecular 
 Hypothesis. 
 
 Definitions of a Molecule. 
 
 (1) A molecule is the smallest particle of matter which can 
 exist alone and retain the properties of a substance. 
 
 (2) " Molecules are the imaginary units of which bodies are 
 the aggregates." Smith. 
 
 (3) A molecule is the physicist's unit.* 
 
 Application of Kinetic-Molecular Hypothesis to, 
 
 (1) Solids, 
 
 (2) Liquids, 
 
 (3) Gases. 
 
 Boyle's Law. 
 
 Charles' Law (Dalton Gay-Lussac Charles). 
 
 Avogadro's Rule. 
 
 Gas Law (PV = RT). 
 
 Critical Phenomena. Experiments with " critical" tubes. 
 
 (a) Critical Temperature (H, 2 3 8C.), (O, n8C.), (Ether, 
 
 .i95C.). 
 
 (b) Critical Pressure (H, 20 atmos.), (O, 51 atmos.), (Ether, 
 
 36 atmos.). 
 
 (c) Critical Point. 
 
 (d) Critical Volume. 
 
 Combining Volumes. Recall Exp. and Law as discussed under 
 "Quantitative Relationships." Two volumes of hydrogen, say 40 cm. 3 
 combine with one volume, 20 cm. 3 of oxygen to form two volumes, 40 cm. 3 
 of steam (water). 
 
 n n + n n n 
 
 2 Vol. Hydrogen, i Vol. Oyxgen. 2 Vol. Steam, 
 (i molecule of oxygen must divide.) 
 
 n + n n n 
 
 Hydrogen. Chlorine. Hydrogen Chloride. 
 
 (Both molecules must divide.) 
 
 * The modern physicist will probably insist that the " electron " is the unit. 
 4 AO 
 
$0 EXPERIMENTAL CHEMISTRY. 
 
 Statement and Historical Development of Atomic Theory. 
 
 Definitions of an Atom. 
 
 (1) An Atom is the smallest particle of matter which can 
 participate in a chemical reaction. 
 
 (2) Atoms are the imaginary units of which molecules are the 
 aggregate. 
 
 (3) An atom is the chemist's unit. 
 
 (4) An atom always represents a distinct variety of matter. 
 
 Statement of J. J. Thomson's Electronic Theory of Matter. See 
 Thomson's " Electricity and Matter." 
 
 Application of Atomic Theory to, 
 
 (1) Chemical Reaction. 
 
 (2) Quantitative Relationships. 
 
 Atomic Weight " is the number which represents the smallest mass oj an 
 element which is known to participate in a chemical change, relative to the 
 smallest weight oj hydrogen which can so junction" Newth; or, "it is 
 the unit oj weight actually used in expressing the proportions oj each ele- 
 ment in all oj its compounds" Smith. 
 
 Note. Hydrogen has been practically abandoned in favor of oxygen 
 as the standard of atomic weights. 
 
 Nomenclature of Atomic Theory. 
 Symbols (Hg, O, H, etc.). 
 Molecular Formulae. 
 
 (a) Simple Molecules (H 2 , O 2 , O 3 , Hg, As 4 , S 8 ). 
 (6) Compound Molecules (FeS, HgO, H 2 O, HC1, MgO). 
 
 Molecular Weight is equal to the sum oj the atomic weights oj all oj the 
 atoms oj which the molecule is composed. 
 
 DETERMINATION OF MOLECULAR FORMULA AND EQUATIONS. 
 Equations. 
 
 (a) Matter or Mass Equations. 
 
 EXAMPLES: 
 
 H 2 + O -* H 2 O. 
 H 2 + C1 2 2HC1. 
 Mg + O MgO. 
 
 (b) Energy Equations. 
 
 EXAMPLES: 
 
 H 2 + O H 2 O + 68,000 cal. 
 
 H 2 + C1 2 2HC1 + 44,000 cal. 
 
 C + O 2 CO 2 + 97,000 cal. 
 
 H 2 + I 2 * 2HI 12,000 cal. 
 
 Hg + O HgO + 22,000 cal. 
 
OUTLINES. 51 
 
 When a body or substance, by virtue of its position or its motion or 
 its state, is capable of doing work it is said to possess energy ; if resistance 
 is overcome, work is done. Atoms do not exist as individuals except 
 in a few instances (Hg, Na, K, Zn, Cd, He, A, Kr, Ne, X); they are 
 united to form molecules, simple or compound, of varying degrees of 
 stability. The stability of this equilibrium is evidently determined by 
 some force operative between the individual atoms ; for want of a better 
 term we designate this force "chemism" or " chemical affinity" (see 
 Note on Energetics of Chemistry). To return to our original line of 
 thought, it is evident that atoms possess energy, because in uniting they 
 are capable of performing work. Atoms are said to possess "chemical 
 energy" (see Note on Energetics of Chemistry) in virtue of their tendency 
 to combine with other atoms. 
 
 It is probable that the energy liberated when atoms combine is due to 
 a transformation of all or a part of this chemical energy into jree energy, 
 i.e., energy which is available for the performance of work. If certain 
 kinds of atoms have a great chemical affinity for one another they, that 
 is the system which they represent, possesses a large amount of chemical 
 energy which will probably be converted into a correspondingly large 
 amount of free energy. Molecules formed under such conditions possess 
 much smaller quantities of energy than their constituent atoms originally 
 possessed, consequently they are stable. Again, two or more atoms 
 may possess so very little chemical affinity for each other that it will be 
 necessary to contribute energy to secure their combination; molecules 
 formed under these conditions will have a tendency to convert this 
 acquired energy into free energy and therefore, will be unstable. 
 
 The atoms of the elements show great variations in respect to the 
 chemical affinity which they manifest toward one another; the same is 
 true as regards the interaction of groups of atoms. 
 
 Valence, or Quantivalence is the atom -fixing or replacing power of an 
 atom of an element in terms of the hydrogen atom which is considered as 
 having a valence of one. Those atoms which have a valence of one, 
 two, three, etc., are said to be respectively univalent, bivalent, trivalent, etc. 
 
 Multiple Valence. The valence of an elementary atom "is not an 
 invariable property oj the atonij but each of the observed differences as to its 
 valence is an invariable property of some particular condition oj the atom 
 dependent on the circumstances in which it is placed." 
 
 Formulae. 
 
 Empirical, Ex., H 2 O, C 2 H 4 O 2 . 
 
 Rational, Ex., HC 2 H 3 O 2 . 
 
 Dualistic, Ex., H 2 O.SO 3 , Na 2 O.SO 3 , 
 
 H \ H\ H] 
 Typical, Cl / ' H / U ' H N. 
 
 H 
 
52 EXPERIMENTAL CHEMISTRY. 
 
 H H H-0 O 
 
 / / \ / 
 
 Structural or Graphic, O , N-H, Na OH, S 
 
 \ \ / \ 
 
 H H H-O O 
 
 Space (formulae of more than one dimension, to represent the con- 
 figuration of a molecule). 
 
 Units of Measurement. 
 
 Length Cm., mm. 
 Volume Cm. 3 , liter. 
 
 Weight. 
 
 (a) Gram. 
 
 (b) Gram-molecule (mole). 
 
 (c) Formula- Weight. 
 
 Temperature Centigrade Scale. 
 
 (a) Boiling Point (iooC.). 
 
 (b) Freezing Point (o C.). 
 
 (c) Absolute Zero ( 273 C.). 
 
 Heat Units Calorie. 
 
 (a) Small Calorie. 
 
 (b) Large Calorie. 
 
 (c) Ostwald Calorie. 
 
 Calculations in Chemistry (Stoichiometry). Problems. 
 
 Mg + O MgO. 
 
 24 1 6 40. 
 KC1O 3 -*KC1 + 30. 
 122.5 74-5 48. 
 
 DETERMINATION OF MOLECULAR AND ATOMIC WEIGHTS. 
 
 Methods. 
 
 (1) Volumetric. (Molar-volume 22.4 1.). 
 
 (2) Chemical. 
 
 (3) Specific Heats Atomic Heat. 
 
 Dulong and Petit's Law Neumann's Rule. 
 
 (4) Isomorphism. 
 
 (5) Elevation of Boiling Point. 
 
 (6) Depression of Freezing Point. 
 
 References. 
 
 Atomic Theory. Any large text-book in chemistry. Theoretical 
 Chemistry. Remsen. Heroes of Science, Chemists. Muir. Atomic 
 Theory. Wurtz. Alembic Club Reprints, No. II. 
 
 Avogadro's Rule. Chemical Theory. Dobbin and Walker. Out- 
 lines of General Chemistry. Ostwald. 
 
OUTLINES. 53 
 
 Critical Phenomena. Introduction to Physical Chemistry. Walker. 
 Any text-book on physical chemistry. 
 
 Valence. Theoretical Chemistry. Remsen. Theoretical Chemistry. 
 Nernst. 
 
 Specific Heats (Dulong and Petit's Law). History of Chemistry. 
 Ladenberg. Outlines of Theoretical Chemistry. Meyer. 
 
 Elevation of Boiling Point, etc. Physical Chemistry. Walker. 
 Physical Chemistry. Jones. Theoretical Chemistry. Nernst. Mod- 
 ern Theory of Solution. Jones. See later work under the subject 
 of " Solutions." 
 
 PROBLEMS. 
 
 Note. See At. Wt. tables in Appendix. 
 
 i. What is the molecular weight of HgO, H 2 O, FeS, MgO, HC1, 
 H 2 SO 4 , HNO 3 , NH 4 OH, NaOH, KOH? 
 
 2. Mg -f O MgO. 5 grams of magnesium will combine with 
 how many grams of oxygen ? Now many grams of magnesium oxide will 
 be formed? 
 
 3. How many grams of hydrogen in 10 grams of H 2 O ? 
 
 4. How many grams of oxygen can be prepared from 10 grams of 
 mercuric oxide? HgO > Hg + O. 
 
 5. Fe + S FeS. If 10 grams each of sulphur and iron are placed 
 in the test tube and chemical action results with a transfer of all of the 
 iron into iron sulphide, how much sulphur remains in the uncombined 
 condition ? 
 
 6. How much oxygen can be prepared from two gram-molecules of 
 potassium chlorate? KC1O 3 KC1 + 30. 
 
 7. A certain weight of KC1O 3 , was heated until completely decom- 
 posed as per above equation. The residue weighed 10.123 gm. What 
 was the original weight of KC1O 3 , and how much oxygen was evolved? 
 How many liters of oxygen were liberated if the standard weight of a 
 1. of oxygen is 1.4 gm. ? 
 
 8. At 760 mm. and o C. how many liters of gas in 2 gm. (molecular 
 wt.) of hydrogen? In 16 gm. (.5 mol. wt.) of oxygen? In 32 gm. 
 (mol. wt.) of oxygen? What is the molar volume? 
 
 9. (a) 5.6 1. of a gaseous substance weights 4.5 gm. approximately, 
 at standard conditions ; what is its molecular weight ? 
 
 (b) If | by weight of this substance is hydrogen and f oxygen, 
 what is its formula? 
 
 10. (a) If the specific heat of Ag (silver) is approximately .057, 
 what is its atomic weight ? Of Fe, if sp.ht. is .112 ? 
 
 (b) If the sp.ht. of Ag Cl (silver chloride) is .089, prove that the 
 atomic weight of Cl is nearer 35 than 70. What is the atomic heat of 
 chlorine ? 
 
 n. Write the structural formula for NH 3 , H 2 O, MgO, FeS, HC1, and 
 H 2 SO 4 , if in this latter compound (sulphuric acid) sulphur has a valency 
 of 6. 
 
CHAPTER VIII. 
 OXYGEN. 
 
 Symbol O. At. Wt. 16 (15.9). 
 
 Oxygen is an elementary substance belonging to the class previously 
 designated as " non-metals." It was first discovered by Priestly, who 
 in 1774 prepared it "by heating 'red precipitate' (red oxide of mercury) 
 in the focus of a burning glass exposed to the sun's rays." It is quite 
 probable that Scheele, a Swedish chemist, had previously obtained it, 
 but he delayed publication of his results until 1777. Lavoisier, a French 
 chemist, named the new element "oxygen" which signifies acid-pro- 
 ducer. He believed that all acids owe their characteristic properties to 
 the presence of oxygen. Cavendish pointed out that many substances 
 which do not contain oxygen possess acid properties. We now know 
 that Lavoisier's view was incorrect. 
 
 Occurrence. Oxygen is the most abundant element upon our planet; 
 it forms about 47 per cent, of the solid portion of the earth and consti- 
 tutes 20 per cent, of the earth's atmosphere. Including its occurrence 
 in the ocean, vegetable and animal forms, etc., it constitutes 50 per 
 cent, of the total substance of the globe. It occurs most abundantly 
 in the free state in the air. 
 
 Physical Properties. .Oxygen is a colorless, odorless and tasteless 
 gas; it is slightly heavier than the air; its density as compared with air 
 which is considered as i (physical standard), is 1.105. One liter of 
 oxygen at standard conditions weighs 1.429 gm. It is slightly soluble in 
 water, 100 volumes dissolving 4. i volumes of oxygen at 5 C. Its critical 
 temperature is 118 C. ; its critical pressure, 50 atmospheres. Liquid 
 oxygen boils at 182 C. at 740 mm. pressure (Dewar and Fleming); 
 it has a pale-blue color; its sp. gr. at 182.5 C. is 1.13. 
 
 Chemical Properties. By chemical properties is meant "reaction 
 properties" Oxygen possesses an almost universal affinity, i.e., "it 
 forms chemical compounds, called oxides, with all other elements ex- 
 cepting fluorine and bromine and it will combine with the latter element 
 provided some metal is also a constituent of the compound." Freer. 
 Much heat is liberated during many of these oxidations, i.e., formation of 
 oxides. This indicates that oxygen possesses considerable chemical 
 energy. Substances which undergo oxidation also possess a certain 
 amount of chemical energy in the presence of oxygen, and when they 
 combine with the latter element this energy is converted into heat. 
 Oxygen is ordinarily bivalent. 
 
 54 
 
OXYGEN. 55 
 
 Experiment I. Sources of Oxygen. 
 
 (a) Place a small quantity of mercuric oxide (HgO) in a small test 
 tube. Note its color. Heat bottom of tube strongly and while doing 
 so, introduce a glowing splinter of wood. Results? Continue to heat 
 until sides of t.t. show a metallic lustre. Explain. Write the equation 
 for this reaction. When the test tube is set aside and permitted to cool, 
 does the powder in the bottom of the tube regain its original color? 
 Explain. Write the equation for this latter reaction. Is it a reversible 
 reaction ? 
 
 (b) Repeat (a) using separately small quantities of MnO 2 , CuO, 
 KC1O 3 , and PbO 2 . 
 
 (c) Place a piece of sodium peroxide (Na 2 O 2 ), about the size of a pea, 
 in a test tube half filled with water. Test the evolved gas as in a (?). 
 Equation ? 
 
 Hg + O > HgO + 30,600 cal. 
 
 How many calories of heat would have to be supplied to a gram- 
 molecular weight of HgO to decompose it into Hg and 6 ? 
 
 Experiment II. Laboratory Source of Oxygen. 
 
 Thoroughly mix 10 grams of potassium chlorate (KC1O 3 ) with exactly 
 8 grams of MnO 2 (manganese di- or per-oxide) ;by means of a folded piece 
 of paper introduce the mixture into the large test tube or retort (Fig. 8) 
 
 which you will find in the drawer; close the tube with a rubber cork fitted 
 with delivery tube; clamp the tube to ring stand at an angle of about 
 45 so that delivery tube will dip beneath the surface of the water in the 
 pneumatic trough ; heat lower end of tube cautiously, and fill six bottles 
 with the gas by displacement of water; place bottles in upright position 
 and cover with wet paper. Observe color, taste and odor of gas as it is 
 
56 EXPERIMENTAL CHEMISTRY. 
 
 evolved; remove delivery tube from pneumatic trough; when tube is cold, 
 add water to contents and allow to soak for several days if necessary 
 to remove mixture ; see Exp. VII. Write equation for reaction, remember- 
 ing that MnO 2 is not altered by the reaction. Proceed with following 
 experiments. 
 
 PROPERTIES OF OXYGEN. "KINDLING TEMPERATURE." 
 OXIDATION OXIDES. 
 
 Experiment III. Oxidation of Carbon. 
 
 (a) Drop a small piece of charcoal into one of the bottles containing 
 oxygen; immediately recover bottle. Is there any evidence of chemical 
 action? Heat to redness another small piece of charcoal; by means 
 of a pair of forceps, deflagrating spoon or a piece of wire, introduce 
 glowing charcoal into above bottle of gas. Results ? 
 
 (b) Pour a little clear lime-water (Ca(OH) 2 ) into another bottle of 
 the gas; place your hand over mouth of bottle and shake; does the water 
 change in color? Now introduce a piece of glowing charcoal as per 
 above experiment; when charcoal ceases to burn, shake bottle as before. 
 Is there any change in the color of the water? Write equations repre- 
 senting above reactions. If 97,000 calories are liberated by the form- 
 ation of one gram-molecule of CO 2 from free O and free C, write the 
 "energy equation" for the reaction. 
 
 The "kindling temperature" of a substance is the temperature to 
 which it must be raised before it will undergo combustion. It is definite 
 for a particular substance but varies greatly for different substances. 
 
 Experiment IV. Oxidation of Phosphorus. 
 
 Yellow phosphorus is usually moulded into small cylinders. Place 
 one of these pieces under water and then by means of a knife cut off 
 a piece about the size of a small pea. Never handle P with the hands, 
 as it readily catches fire. Cover the bottom of one of the bottles of gas 
 with water; test water with litmus paper (blue); place phosphorus in a 
 deflagrating spoon (Fig. 9) and ignite it, then lower it into bottle of gas. 
 Results ? Test water in bottle with blue litmus paper. Write equations 
 for reaction. Was energy set free by the reaction ? 
 
 Experiment V. Oxidation of Sulphur. 
 
 By means of a deflagrating spoon, lower a little ignited sulphur into 
 a bottle of gas; the bottom should be covered with water and tested with 
 litmus paper before and after the burning of the sulphur. Results? 
 Was energy liberated? Equations? 
 
 Experiment VI. Oxidation of Iron. 
 
 Lower the unwound end of a piece of "picture cord" (iron) into a 
 bottle of the gas for a few seconds. Is there any manifestation of chem- 
 
OXYGEN. 57 
 
 ical action? Remove 'the wire and dip it into the burning sulphur in 
 the deflagrating spoon; if the sulphur on the wire is burning, lower wire 
 into bottle (Fig. 10) again. Results ? The bottom of the bottle should be 
 covered previously with water to protect it from the molten iron. Write 
 equation. Why was it necessary to burn sulphur on the end of the 
 wire ? Was energy set free by the oxidation of the iron ? 
 
 The greatest diversity of characteristics exists among these oxides, 
 which for purposes of convenience may be classed either as metallic or 
 non-metallic oxides. Many of the elements are capable of forming a 
 series of oxides; nitrogen, for example, forms five different compounds 
 with oxygen. When more than one oxide is formed, the suffix ous or ic 
 is attached to the name of the metal accordingly as the oxide, relative to 
 the metal, contains a lesser or greater amount of oxygen, i.e., these 
 endings indicate the relative degrees of oxidation which the metals have 
 
 FIG. 9. FIG. 10. 
 
 undergone, thus iron (ferrum) forms two oxides, ferrous oxide (FeO), 
 and ferric oxide (Fe 2 O 3 ) ; mercury possesses the same.property of forming 
 two oxides, mercurous (Hg 2 O) and mercuric (HgO) oxides; copper 
 forms cuprous oxide (Cu 2 O), and cupric oxide (CuO). 
 
 Inasmuch as oxygen is ordinarily bivalent, it is obvious that a metal 
 frequently shows a varying valency as represented by its ous and ic 
 compounds. 
 
 In the preceding experiments involving "oxidation," it was readily 
 noted that the reactions were exothermic, i.e., the "heat of formation" 
 of the respective oxides was positive; it follows that the energy (usually 
 heat) necessary to decompose ("heat of decomposition") an oxide is 
 equal to the energy (usually light and heat) liberated during its formation. 
 
 This may be made clear by the use of an analogy. As mentioned 
 previously, substances capable of directly combining, possess a definite 
 amount of chemical energy which may be likened unto the potential 
 energy which a stone possesses when it is raised above the ground; when 
 the substances combine, all or portions of the chemical energy of the 
 factors is transformed into say heat and light energy, just as the potential 
 energy of the stone is converted into kinetic energy by permitting the 
 stone to fall ; now it will require just as much energy to raise the stone to 
 
58 EXPERIMENTAL CHEMISTRY. 
 
 its original elevated position as was freed during its descent; just so 
 with chemical compounds, the "heat of decomposition" is necessarily 
 equal to the "heat of formation." If combination is attended with the 
 evolution of much energy (heat and light) the oxides will be stable, 
 and vice versa. 
 
 Although the "heat of combustion" of carbon in oxygen with forma- 
 tion of CO, can not be directly measured, it can be calculated from 
 calorimetric measurements indicated by the following equations: 
 
 C + 2O CO 2 + 97,000 cal. 
 
 and CO + O CO 2 + 68,000 cal. 
 
 subtracting, C + O > CO + 29,000 cal. 
 
 Experiment VII. Catalytic Action of Manganese Dioxide in Exp. II. 
 
 This experiment is a continuation of Exp. II. Half fill the test tube 
 with water; heat gently in an endeavor to dissolve contents of tube; be 
 careful to avoid losing any of the solid material; pour the clear super- 
 natant liquid upon a filter; avoid pouring much of the solid matter upon 
 the filter; half fill the tube again with water and heat to boiling; decant 
 the fluid as before upon the filter; repeat above operation four or five 
 times, then pour contents of tube on filter; rinse out tube thoroughly, 
 and wash material upon filter with hot water; spread filter and contents 
 upon wire gauze and heat gently until paper is dry; place paper upon 
 glass plate and scrape black powder into a weighed crucible ; fold filter 
 paper and place it in crucible ; heat crucible with low flame until contents 
 are perfectly dry, and filter paper has ignited and burned to ash ; avoid 
 heating crucible to redness; cool and weigh; ignore weight of ash of 
 filter paper; what is the weight of the black powder? Does this 
 powder resemble MnO 2 ? Do your results justify the conclusion that 
 WnO-j was not altered in the reaction ? Add a few drops of the filtrate to 
 a cubic centimeter of silver nitrate. Results ? Add a few drops of a 
 potassium chlorate solution to a cubic centimeter of silver nitrate. Re- 
 sults ? Did the KC1O 3 undergo a change in Exp. II ? 
 
 2KC1O 3 + 2MnO 2 2KMnO 4 + 2C1 + 2O 
 2KMnO 4 K 2 MnO 4 + MnO 2 + 2O_ 
 K 2 Mn0 4 + 2C1 2 KC1 + MnO 2 + 2!) 
 adding, 2 KC1O 3 + 2MnO 2 2KC1 + 2 MnO 2 + 6O(3O 2 ). 
 
 A substance which alters the speed of a chemical change apparently 
 by its mere presence and contact, without undergoing any permanent 
 change, is called a catalytic agent or simply, a catalyser. It may be either 
 a gas, liquid or solid. Pieces of platinum foil are frequently used to 
 catalyse a reaction. The process itself is called catalysis. 
 
 "To obtain a picture of the way in which a catalyser acts, imagine a 
 wheel-work in which the axles move with great friction, as a result say of 
 the oil having become thick, and which therefore runs down only very 
 
OXYGEN. 59 
 
 slowly. If a little fresh oil be placed on the axles the wheel-work forth- 
 with runs down much more quickly, although the available tension of 
 the spring (which corresponds to the work available from chemical 
 reaction) is in no way altered by the oil. The action of a catalyser may 
 be compared with that of the oil in this respect, and also with respect 
 to the fact that the oil is not used up in acting." Ostwald. 
 
 It should be understood that a catalyser does not initiate a reaction; 
 it is not the cause of a reaction; it merely increases or diminishes the 
 speed of a given chemical reaction. 
 
 Experiment VIII. Preparation and Properties of Ozone (O 3 ). 
 
 Place two or three pieces of yellow phosphorus, about the size of a 
 " playing" marble, in a bottle; add a sufficient quantity of water to cover 
 bottom of bottle, but not any more than will be necessary to half cover 
 the pieces of phosphorus; four or five drops of a solution of potassium 
 dichromate should be added to the water; cork the bottle. Now prepare 
 a piece of test-paper as follows: dip a strip of filter paper into a starch 
 emulsion to which has been previously added a few drops of a solution 
 of potassium iodide (KI); let the strip of paper "drip" for a few minutes, 
 then suspend it from the cork in the bottle so that paper is within a centi- 
 meter of the phosphorus; allow the apparatus to stand for one laboratory 
 period, then examine paper. Note the odor of the gas in the bottle. 
 Results? Equations? The phosphorus should be returned to the 
 supply bottle. 
 
 Note. A good test paper for above experiment is made by dipping 
 a piece of red litmus paper into a solution of KI. When the O 3 acts upon 
 the KI, the iodine is set free and KOH acts upon the red litmus paper 
 turning it blue. Satisfactory results are frequently secured by placing 
 phosphorus in a dry flask which is then closed with a cork. 
 
 A quantitative study of the conversion of oxygen (O 2 ) into ozone 
 (O 3 ) reveals that 3 volumes of the former are required to form 2 volumes 
 of the latter, and that the reaction is endothermic, i.e., absorbs heat 
 energy. 
 
 (1) 2 + O 3 32,900 cal. 
 
 (2) O 3 O 2 + O + 32,900 cal. 
 
 At a temperature of about 275 C. the above reaction (i) is reversed. 
 Ozone is a gas of blue color, odor like dilute chlorine, density 24 (H 2 = i); 
 molecular weight, 47.9; boils at 119 C. Ozone is much more soluble 
 than oxygen in water 100 volumes of water at 12 C. dissolves 50 
 volumes of ozone under a pressure of one atmosphere. The chemical 
 properties of oxygen and ozone are similar, save that the latter is much 
 more active. 
 
 It is evident that oxygen and ozone are identical so far as the kinds 
 
60 EXPERIMENTAL CHEMISTRY. 
 
 of matter of which they are composed is concerned. It seems to be 
 equally as evident that their different properties are due to the fact that 
 they differ in their respective energy contents. When ozone is used 
 as an oxidizer, 32,900 more calories are liberated than when an equal 
 weight of oxygen is used, hence the activity of ozone as an oxidizer. 
 It is to be noted that the heat developed by the passing of ozone into 
 ordinary oxygen is not contained in the ozone, as heat energy but in the 
 form of chemical energy. 
 
 Ozone and oxygen are allotropic modifications of the same element. 
 
 Experiment IX. (Quant. ) Determination of, (a) the Density of Oxygen, 
 (b) the Volume of Oxygen Liberated from a Given Weight of Potassium 
 Chlorate, (c) the Weight of the Potassium Chloride Formed. 
 
 (a) Note. i gram of KC1O 3 liberates about 290 cm. 3 of O 2 . 
 
 Place a gas burette, w r hich is connected with a leveling bulb, in a water 
 jacket; suspend a thermometer in the water jacket; raise leveling bulb 
 until water in burette stands at top of capillary tube on upper end of 
 burette. Clean and dry and weigh a hard- glass test tube; record weight; 
 introduce into the tube .12 grm. of KC1O 3 which has been previously 
 pulverized, and dried in an air bath (100 C.) for one hour. Connect test 
 tube with burette by means of rubber tubing and a piece of capillary 
 glass tubing bent into such form as will permit test tube to be clamped at 
 an angle of about 45 . All joints must be air tight. When making the 
 connections be sure that no water is forced into the glass tubing. Adjust 
 the leveling bulb so that \vater in it and in burette are at the same level. 
 Read burette and thermometer, and record readings. Holding burner 
 in hand gently warm the test tube; oxygen will be liberated and pass 
 over into the burette. Always keep the water in the leveling bulb about 
 3 or 4 cm. below the water in the burette. Continue to heat tube until 
 gas ceases to be evolved, then remove burner and allow the apparatus 
 to assume its original temperature, then bring the water in burette and 
 bulb to same level. Read burette, thermometer and barometer, and 
 record readings. From tables in Appendix get vapor tension (a) of 
 water at the temperature of the water jacket. Reduce the volume 
 of oxygen to standard conditions. 
 
 Vol. x 273 x (Bar. P. a) 
 
 (273 + t) x 760 
 
 Remove the test tube; wipe and weigh it; what is the weight of the 
 KC1? The difference between the weight and the original weight of 
 the KC1O 3 is the weight of the oxygen evolved. Calculate the weight 
 of a cm. 3 of oxygen at standard conditions. The weight of a liter? 
 Calculate the weight and volume of oxygen which should have been 
 liberated from the KC1O 3 on a purely theoretical basis. 
 
OXYGEN. 
 
 6T 
 
 Experiment X. (Quant.) Optional. Determination of the Percentage of 
 Oxygen in the Air. 
 
 Note. The general theory of this experiment is to isolate a given 
 volume of air under known conditions, then bring this volume of air into 
 intimate contact with a substance which will absorb the oxygen, after 
 which the volume of air is measured under the original conditions, 
 and its loss in volume noted. 
 
 Use the measuring apparatus assembled for Exp. IX or Hempel 
 burette (Fig. n). Attach a piece of capillary rubber tubing to the 
 upper end of the gas burette; the tubing should be about 7 cm. long 
 and firmly wired to the burette; place a Mohr pinchcock upon the 
 rubber tubing close to the end of the burette; open pinchcock; raise 
 leveling bulb until water in burette stands at top of capillary tubing, 
 then lower bulb until about 40 cm. 3 of air have been drawn into 
 burette; close pinchcock; allow gas to stand for five minutes; bring 
 water in burette and bulb to the same level; read thermometer sus- 
 
 FIG. ii. Hempel Burette. 
 (Smith and Keller.) 
 
 FIG. 12. Hempel Compound Pipette. 
 (Smith and Keller.) 
 
 pended in water jacket, read barometer, read burette, and make a record 
 of readings; now a compound gas pipette (any desirable form, Fig. 12) 
 containing an alkaline solution of pyrogallol,* should be connected 
 with the gas burette by means of a short piece of capillary glass tubing, 
 the pipette should be closed by rubber tubing and a pinchcock; when the 
 pipette and burette are properly connected slightly raise the leveling 
 bulb, then open both pinchcocks and by gradually raising the bulb the 
 air will be forced over into the pipette; raise bulb until water in burette 
 stands at top of rubber tubing on burette; close both pinchcocks; dis- 
 connect pipette and burette; shake pipette for 20-25 m i n -5 connect 
 pipette with burette; gradually lower bulb, at the same time open the 
 pinchcocks; continue to lower bulb until all the gas is drawn from pipette 
 
 *The alkaline " pyro " solution is prepared as suggested by Professor Baxter of 
 Harvard, as follows 
 
 One gram of pyrogallol to 4 cm. 3 of water; i gram of KOH to i cm. 3 of water. The 
 solution prepared in above ratio must be mixed out of contact with the air; the pyro 
 solution is first introduced into the pipette after which the KOH solution follows; 
 shake until solutions are thoroughly mixed. 
 
62 EXPERIMENTAL CHEMISTRY. 
 
 and the "pyro" solution is drawn up to the top of rubber tubing on 
 pipette; close pinchcocks, be sure that both are closed tightly; allow 
 gas in burette to stand for 5 min.; bring water in burette and bulb to 
 same level; read burette, thermometer, barometer and record readings. 
 Repeat the operation until no further decrease in volume occurs. Reduce 
 both the original and final volumes of air to standard conditions. What 
 is the total change in volume? 'Calculate the percentage of oxygen in 
 the air from your measurements. 
 
 A principle of great importance, especially in connection with the 
 measurement of the volume of gases, is incorporated in Dalton's Law 
 of Partial Pressures: 
 
 The pressure exerted by a mixture oj a gas and a vapor, of two vapors, 
 or oj two gases, is equal to the sum oj the pressures which each would exert 
 if it alone occupied the whole space afforded to the mixture. 
 
 This law is not absolutely exact. It may be stated in another form: 
 
 The volumes oj two or more gases in a volume oj the mixture are pro- 
 portional to their respective pressures. 
 
 Another of Dalton's Laws is: 
 
 The pressure exerted by, and the quantity oj a vapor which saturates a 
 given space are the same for the same temperature whether this space is 
 filled by a gas or is a vacuum. 
 
 PROBLEMS. 
 
 i. How many grams of sulphur will exactly combine with 32 grm. 
 of oxygen to produce SO 2 ? 
 
 2. What weight of oxygen could be obtained from 10 grm. of KC1O 3 
 if the latter contains 10 per cent, of an impurity? Now many liters, if 
 a liter of oxygen weighs 1.428 grm. ? 
 
 3. What is the "heat of formation" of SO 2 if 2 grm. of sulphur in 
 burning to sulphur dioxide develops 4440 calories? Write the "energy" 
 equation. 
 
 4. The specific heat of copper is .092. How many calories of heat 
 are liberated when 100 grm. of copper cool from 75 C. to 50 C. ? 
 
 5. Is the heat liberated during formation of an oxide, necessarily 
 an accurate measure of the chemical energies (affinity) of the reacting 
 substances ? Explain. 
 
 6. Reduce 200 cm.s of gas at 20 C. and 730 mm. to o and 760 mm. 
 
 7. Reduce 45 cm. 3 of gas at 20 C. and 770 mm. to oand 760 mm. 
 
 8. Reduce 40 cm. 3 of gas collected over water at 20 C. and a baro- 
 metric reading of 730 mm. to o C. and 760 mm. 
 
 9. A gas globe when full of air weighed 55.06 grm.; full of water at 
 
OXYGEN. 63 
 
 20 C. it weighed 309.66 gram. The globe was " exhausted " (air removed) 
 and carefully weighed by means of a counterpoise; it was then filled with 
 a gas at 19.8 C. and 761.4 mm., when it weighed .469 grm. more than 
 when exhausted. What is the weight of a liter of this gas at standard 
 conditions ? 
 
 10. What is the molecular weight of the gas used in above experi- 
 ment? 
 
CHAPTER IX. 
 HYDROGEN. 
 
 Symbol H. At. Wt., 1.0075. 
 
 Hydrogen is a light colorless gas, first described as a form of air 
 (phlogiston) by Cavendish, in 1776, although it had undoubtedly been 
 observed previously by Paracelsus and Boyle. Lavoisier gave the ele- 
 ment the name, hydrogene (to produce water). It is univalent. 
 
 Experiment I. Methods for Preparing Hydrogen. 
 
 (a) Electrolysis of water (H 2 O). Recall, or repeat this experiment 
 as previously performed. At which pole was the hydrogen liberated? 
 What test for hydrogen was used ? 
 
 (b) I. Decomposition of water by means of metals. Wrap a piece 
 of sodium or potassium about the size of a pea in a small piece of filter 
 paper; place it in a small wire cage and dip it under the surface of the water 
 in the pneumatic trough (Fig. 13); hydrogen gas will be evolved, which 
 collect in a test tube by displacement of water; when action has ceased 
 place your thumb over the mouth of tube and hold tube in upright 
 position; place a glowing splinter in the mouth of tube for a second. 
 Results? Bring a lighted match to the mouth of the tube. Results? 
 Does the hydrogen support combustion so far as the glowing match is 
 concerned ? When raised to a given temperature does it combine with a 
 component of the air? 
 
 H 2 + K->KOH + H 
 Water Potassium 
 
 Hydroxide, 
 Potassium, Hydrogen. 
 
 H 2 O + Na NaOH + H. 
 
 (b) II. In the above experiment all of the hydrogen of the molecule 
 of water was not displaced by the metal; this may be accomplished by 
 another process. Pulverize and mix thoroughly and quickly about i 
 grm. of sodium hydroxide (NaOH) and 2 grm. of zinc dust; introduce 
 mixture into a hard glass test tube; heat tube and test for hydrogen gas. 
 Results ? 
 
 Zn + 2NaOH Zn(ONa) 2 + 2 H. 
 
 (b) III. Decomposition of water by magnesium. Shake a gram of 
 magnesium powder into the bottom of a test tube and add 10 cm. 3 of water; 
 
 64 
 
HYDROGEN. 65 
 
 add a few cm. 3 of magnesium chloride (this does not take any part in the 
 reaction, merely dissolving the magnesium oxide which is formed on the 
 surface of the metal). Heat the tube and its contents; test for hydrogen 
 gas. Results ? Equation ? 
 
 (c) Displacement of Hydrogen from Acids by Metals. 
 
 (c) I. To a small piece of each of the following metals, placed in sep- 
 arate test tubes add 5-10 cm.3 of hydrochloric acid, HC1 (shelf-reagent); 
 granulated zinc, magnesium (ribbon), iron (filings). Test each tube 
 for the presence of hydrogen. Notice whether the test tubes become warm 
 as the reaction continues. What is the effect of heating the tubes ? Re- 
 cord all data. Assume that all of the above metals are bivalent; repre- 
 sent the reactions by "energy equations" using the word "heat" instead 
 of calories. Does the final "system" contain more or less energy than 
 the original one? Why? 
 
 FIG. 13. 
 
 (c) II. Repeat (c) I, using cold dilute sulphuric acid (H 2 SO 4 ) instead 
 of HC1. This dilute acid may be prepared by pouring i volume of con- 
 centrated H 2 SO 4 (shelf-reagent) into 4 volumes of water; allow mixture 
 to cool. Add a few drops of copper sulphate solution to test tubes con- 
 taining iron and zinc respectively. What is the effect ? Similar results 
 are secured by placing a piece of platinum foil or wire in contact with 
 the zinc. This may be verified by the student by using a separate test 
 tube containing zinc and dilute H 2 SO 4 . After the zinc has dissolved 
 completely, filter the solution which remains; evaporate the filtrate until 
 a thin film appears on the surface when the solution cools. Are crystals 
 deposited ? Does it resemble zinc or H 2 SO 4 ? Repeat this process using 
 the iron solution. Write equations as in (c) I. Can you explain the 
 action of the platinum or the copper sulphate ? Did the platinum lose its 
 identity ? 
 
66 EXPERIMENTAL CHEMISTRY. 
 
 (c) III. Effect of Surface on Speed of Reaction. 
 
 Try the action of dilute H 2 SO 4 on zinc dust. A small quantity of the 
 latter placed in the bottom of t.t. will be sufficient. Compare results 
 with action of the acid on "granulated" zinc. Explain. 
 
 (c) IV. Compare the interaction of acetic acid (HC 2 H 3 O 2 ) and zinc 
 with foregoing acids and zinc. Repeat above using concentrated H 2 SO 4 . 
 
 The interaction of certain metals with water is identical in principle 
 with the interaction of certain acids and metals; the metallic substance 
 hydrogen, is attached in both cases to a negative element or radical, 
 thus, H.OH, H.C1, H 2 .SO 4 ; the hydrogen (metal) is then displaced by 
 
 FIG. 14. 
 
 certain other metals, (not all metals, for ex. Au., Ag., Pt.) which possess 
 a greater chemical energy when in contact with the above mentioned 
 negative groups. A reaction then occurs whereby a new system is es- 
 tablished which contains less energy than the original one. 
 
 Zn + H 2 SO 4 ZnSO 4 + 2H. 
 
 The " system" represented by the bodies to the lejt of the arrow in the 
 above chemical equation obviously contains more chemical energy than 
 the system represented by those to the right, as evidenced by the fact that 
 heat was given out during the reaction. 
 
 Experiment II. Laboratory Source of Hydrogen. 
 
 Assemble a gas bottle (generating flask); use the heavy glass flask 
 (Fig. 14) of about 500 cm. 3 capacity; fit flask with a rubber cork perforated 
 with two holes; through one of the holes pass a thistle tube; in the other 
 fit a delivery tube with its arms bent at 90; connect delivery tube of 
 generator with a washing bottle ; half-filled with concentrated H 2 SO 4 ; 
 washing bottle should be provided with delivery tube which dips beneath 
 surface of water in pneumatic trough. Place enough granulated zinc 
 
HYDROGEN. 
 
 6 7 
 
 in the flask to cover the bottom completely; add dilute H 2 SO 4 through 
 thistle tube (i of acid to 4 of H 2 O) as needed to secure a rapid evolution 
 of gas; be sure that end of thistle tube dips beneath the liquid in flask. 
 If gas is not evolved rapidly add a few drops of copper sulphate solution 
 through thistle tube. Collect the gas in a test tube by displacement of water 
 until the hydrogen burns quietly in the test tube when a flame is applied 
 
 to it, then collect several test tubes of the 
 gas and proceed with a study of the proper- 
 ties of hydrogen. 
 
 Note. The test tubes containing the 
 gas must remain inverted with mouths 
 under water. Proceed with Experiments 
 IV and V. 
 
 Experiment III. Properties of 
 Hydrogen. 
 
 (a) Remove one of the test tubes con- 
 taining hydrogen from the pneumatic 
 trough and observe whether the gas has 
 any color, odor, etc. 
 
 (b) Allow one of the test tubes filled 
 with hydrogen to stand mouth upward 
 and uncovered for a minute. Apply a 
 flame. Results? Explain. Is the gas 
 apparently lighter or heavier than air? 
 
 (c) Place your thumb over the mouth 
 of a tube of gas; hold it vertically, mouth 
 upward; remove the thumb and immedi- 
 ately bring the mouth of a test tube con- 
 taining air, down over the mouth of the 
 tube of hydrogen, in such a manner as to 
 prevent the gas escaping from the two 
 tubes, but will permit them to "diffuse" 
 into one another; hold them in this posi- 
 tion for i min. ; apply a flame to the 
 
 mouth of each tube. Results ? Explain. 
 
 Tabulate the properties of hydrogen as revealed by above experiments. 
 
 Experiment IV. Effusion of Hydrogen Through a Porous Medium. 
 
 As relatively large quantities of hydrogen are required for this experi- 
 ment (Fig. 15), it can be performed to advantage by the instructor or 
 assistant. 
 
 Law of Effusion, of Graham and Bunsen. The velocities oj effusions 
 of gases are inversely proportional to the square roots oj the densities. 
 Explain on basis of " Kinetic Theory of Gases." 
 
 FIG. 
 
68 EXPERIMENTAL CHEMISTRY. 
 
 Experiment V. Synthesis of Water by Burning Hydrogen in Air. 
 
 Draw a piece of hard glass tubing out into the form of a fine jet. At- 
 tach this jet to the delivery tube of the washing bottle (drying agent in 
 this case); permit jet to dip beneath surface of water in pneumatic trough; 
 fill a test tube with the gas by the usual method; raise jet out of the water; 
 apply flame to test tube of gas; quickly pass the tube over the jet and 
 immediately withdraw it; if gas is pure it burns comparatively slowly so 
 that when tube is passed over jet, the gas issuing from latter is ignited; 
 if gas is impure, i.e., diluted with oxygen, it explodes when flame is ap- 
 plied to test tube. Always light the gas jet by above method. Place over 
 the burning jet a cold dry bottle or let flame impinge upon a cold dry 
 glass plate (Fig 16). Does water form upon the sides of the glass? Is 
 the combination of hydrogen and oxygen to form water an endo- or 
 
 FIG. 1 6. (Smith and Keller). 
 
 exothermic reaction? Your reasons? Explain. Equation? When a 
 lamp in a cold room is first lighted moisture appears on inside of chimney. 
 Explain. When bottles of varying sizes are placed over a burning jet, 
 a musical note is frequently emitted; it has been called the "singing 
 flame." Explain. 
 
 H 2 + O H 2 O, Aq + 68,400 cals. 
 Experiment VI. Hydrogen and Air Form an Explosive Mixture. 
 
 The Hydrogen Cannon. Into a small tin vessel closed at one end 
 with a cork, and containing a small opening at the other end, introduce 
 a stream of hydrogen for a few seconds; apply a flame to the small open- 
 ing after " pointing" cork toward the ceiling. Explain. A mixture of 
 hydrogen and oxygen gases is called "detonating gas." Soap-bubbles 
 filled with this gas will explode when a flame is applied, with a report like 
 a gun. 
 
 Experiment VII. Occlusion of Hydrogen. Catalytic Action of 
 "Spongy Platinum." 
 
 Prepare a piece of " platinized asbestos" by dipping a small piece of 
 " sheet asbestos" in a solution of platinic chloride, and then in an ammo- 
 nium chloride solution; heat asbestos in the hottest portion of the flame 
 of a Bunsen burner or the blast-lamp; finely divided platinum is depos- 
 ited in the asbestos by this process. When asbestos has cooled slightly 
 
HYDROGEN. 69 
 
 hold it by means of forceps or wire so that a current of hydrogen may 
 pass over it. Does the asbestos begin to glow? Is the gas eventually 
 raised to the " kindling temperature " ? Repeat, using natural gas or il- 
 luminating gas. 
 
 "Apparently connected with the catalytic action of platinum is its 
 property of dissolving large quantities of different gases, especially hy- 
 drogen." "The hydrogen thereby increases enormously in reactivity 
 
 " " It must not be supposed that the chemical affinity or the 
 
 chemical potential of the hydrogen is changed; such an assumption, which 
 is certainly very often made, would be a contradiction of the fundamental 
 laws of the theory of energy." "The cause of the changed action of the 
 platinum lies rather in the 'acceleration' of the reactions of hydrogen and is, 
 
 o 
 
 FIG. 17. (Smith and Keller.) 
 
 therefore, a catalytic action. Gaseous hydrogen reacts so slowly at the or- 
 dinary temperature that it appears like an indifferent substance, and 
 from the fact that in the presence of platinum the reaction becomes visible 
 in a short time, while otherwise it would require hours or perhaps years, 
 the view has arisen that there is a change of the chemical potential. "- 
 Ostwald. The finely divided platinum, exposing much surface, is 
 simply a catalytic agent. 
 
 Experiment VIII. Reduction of Cupric Oxide Oxidation of Hydro- 
 gen Synthesis of Water. 
 
 Assemble hydrogen generator; connect it with washing bottle (Fig. 17) 
 which should be half filled with H 2 SO 4 (cone.), that the hydrogen gas shall 
 be dry after bubbling through it; you will find in the drawer a piece of hard 
 glass tubing i cm. in diam. and 18 cm. long; fit into one end a rubber 
 stopper perforated by a single hole; by means of glass tubing connect this 
 end of the tube with the washing bottle the hard glass tube should be in a 
 nearly horizontal position; the entire system may now be flooded with 
 hydrogen; while w r aiting for all of the air to be removed from system, 
 spread i gram, of cupric oxide (CuO) over the bottom of a small porcelain 
 boat which should then be placed in the middle of the horizontal tube; 
 the portion of the tube beneath the boat should be warmed gently at first, 
 and later, quite strongly. Does moisture condense on cold portions of the 
 tube ? Is there any evidence that the CuO has undergone any change ? 
 
70 EXPERIMENTAL CHEMISTRY. 
 
 Remove heat; draw boat out of tube (if water comes in contact with hot 
 tube it will crack) ; notice the color of the powder. Has metallic copper 
 been deposited ? Write equations representing above chemical reactions. 
 
 Note. In above experiment the boat may be dispensed with; place the 
 CuO in the tube. If time permits the reddish-brown powder may be 
 reoxidized by passing a slow current of oxygen over it when heated to 
 redness. Equation ? 
 
 Experiment IX. The Nascent State. Nascent Hydrogen. 
 
 To 10 cm. 3 of a dilute solution of potassium permanganate, KMnO 4 , 
 add an equal volume of dilute sulphuric acid after shaking divide the 
 solution into two parts. To one portion in a test tube add a little zinc dust ; 
 test the evolved gas with a flame. What is the effect of freshly liberated 
 hydrogen, i.e., at instant of liberation or birth, upon the colored solution? 
 
 Pass a stream of hydrogen gas through the other portion of the solution. 
 Results ? Explain. 
 
 Experiment X. Properties of Hydrogen Dioxide (H 2 O 2 ). 
 
 (a) From the "side-shelf" reagents procure 5-10 cm. 3 of H 2 O 2 . 
 Observe its properties, odor, color, reactive properties toward red and 
 blue litmus paper, etc. 
 
 (b) Pour a few drops of H 2 O 2 upon a piece of "test-paper," or better, 
 upon a little starch paste containing a few drops of a KI solution. Re- 
 sults ? Dip a strip of red litmus paper into a dilute solution of potassium 
 iodide (KI); then add two or three drops of H 2 O 2 to litmus paper; explain 
 why litmus paper turns blue. Write equations showing nature of above 
 reactions. 
 
 (c) Heat 5-6 cm. 3 of H 2 O 2 to boiling; test for evolved oxygen; add a 
 pinch of powdered MnO 2 or any powdered metal, preferably the former, 
 to the H 2 O 2 in the test tube; test for oxygen. The MnO 2 acts as a cata- 
 lyser in the reaction :- 
 
 2 H 2 O 2 2H 2 O + "Oi 
 
 (d) Test for chromates. Half fill a test tube with H 2 O; add 10 to 15 
 drops of HC1; add a few drops of potassium dichromate solution to impart 
 a brick red color to solution in test tube; add a layer of ether 2 cm. thick; 
 add two or three drops of H 2 O 2 ; shake. Is the ether colored blue ? This 
 blue color is supposed to be due to presence of perchromic acid (H 2 Cr 2 O 8 ), 
 probably formed as indicated by following equations: 
 
 K 2 Cr 2 O 7 + 2 HC1 . H 2 Cr 2 O 7 + 2KC1 (?) 
 H 2 Cr 2 7 + H 2 2 -> H 2 Cr 2 8 + H 2 O. 
 
 A chromate made strongly acid with HC1 may be used instead of the 
 K 2 Cr 2 O 7 solution. The H 2 O 2 acts as an oxidizer. 
 
 (e) H 2 O 2 as a reducing agent and a bleacher. To a few cm. 3 of a potas- 
 sium permanganate add an equal volume of H 2 SO 4 ; now add sufficient 
 H 2 O 2 to decolorize the solution. 
 
HYDROGEN. 7 1 
 
 KMnO 4 + H 2 SO 4 HMnO 4 + KHSO 4 . 
 2HMnO 4 + 2H 2 SO 4 + sH 2 O 2 2MnSO 4 + 8H 2 O -f 5O 2 . 
 
 (/) Bleaching with H 2 O 2 , a process of oxidation. To a few drops of 
 indigo solution add 5 cm. 3 of H 2 O, then add 4 or 5 cm. 3 of H 2 O 2 ; heat. 
 Results ? 
 
 C 16 H 10 N 2 2 + 2H 2 2 -+ 2C 8 H 5 N0 2 + 2 H 2 O. 
 Experiment XI. Preparation of Hydrogen Dioxide. 
 
 Half fill a test tube with H 2 O; add about 20 drops of dilute HC1 and a 
 sufficient quantity of a potassium dichromate (K 2 Cr 2 O 7 ) to impart a brick- 
 red color to the solution; add ether until there is a layer 2 cm. thick; make 
 a thick paste of BaO 2 and H 2 O; add this paste to the contents of the t.t. ; 
 shake and observe the color of the layer of ether. Compare results with 
 Exp. X (d). Write the equation for the interaction of BaO 2 and HC1. 
 Another method for the preparation of H 2 O 2 is represented by following 
 equation: Decant the solution which is formed as the result of the 
 reaction, 
 
 Ba0 2 + H 2 S0 4 dil. Ba 2 S0 4 + H 2 O 2 , 
 
 into a clean test tube; add ether, and a few drops of a dilute solution of 
 potassium dichromate. 
 
 H 2 + O 2 * H 2 O 2 , Aq + 45,300 cal. 
 H 2 O 2 H 2 O +O + 23,100 cal. 
 H 2 + O H 2 O,Aq + 68,400 cal. 
 
 Experiment XII. (Quant.) Optional. Synthesis of Water. 
 
 Recall or repeat Exp. " Combining Volumes, " under head of " Quanti- 
 tative Relationships." 
 
 Experiment XIII. (Quant.) Hydrogen Equivalent of Zinc, or a 
 Determination of the Volume of Hydrogen Evolved from Sulphuric Acid 
 by a Given Weight of Zinc. 
 
 Assemble gas measuring apparatus described in Exp. X. under " Oxy- 
 gen." Raise leveling bulb until water in gas burette stands at top of 
 capillary on burette; clean and dry a 150 cm. 3 bottle which has a small 
 test tube fused into the bottom (this test tube need not be fused to bottle); 
 fit to bottle a rubber cork perforated with one hole; by means of a short 
 piece of capillary glass tubing connect bottle with top of burette. Clean 
 and dry a piece of pure sheet zinc; weigh a piece of this zinc with great 
 accuracy; the piece must not weigh more than .13 grm. if the capacity of 
 burette is 50 cm. 3 ; carefully wrap a piece of platinum wire around the zinc 
 (the platinum acts as a catalyser); the piece of zinc should how be dropped 
 into the bottle; 25 c.m 3 of pure H 2 SO 4 (i of acid to 4 of H 2 O) is carefully 
 placed in the test tube in the bottle by means of a pipette ; the cork is forced 
 into place and air-tight connections are made with the burette. It is well 
 to place bottle in a water bath; allow the apparatus as assembled to stand 
 
72 EXPERIMENTAL CHEMISTRY. 
 
 for five minutes; level; take temperatures of water bath and water jacket, 
 read burette, read barometer; record all readings; turn bottle until acid is 
 poured out of t.t. upon zinc; gradually lower the leveling bulb as the gas 
 is evolved; after all the zinc has been exhausted and gas is no longer 
 evolved, allow the apparatns to stand for a few minutes; bring the 
 water bath and the water jacket to their respective initial temperatures; 
 bring water in burette and bulb to same level; read burette; reduce vol- 
 ume of gas to o C. and 760 mm. If the weight of i cm. 3 of hydrogen 
 at o C. and 760 mm. is .00009 grm. how many grams of hydrogen will be 
 evolved by 32.7 grams of zinc? 
 
 Calculate both the weight, and the volume (at standard conditions) of 
 hydrogen which should have been evolved from the acid by the weight of 
 zinc used. Compare experimental data with calculated. 
 
 Note. A few drops of PtCl 4 added to the acid could have been sub- 
 stituted for the platinum wire. CuSO 4 is also frequently used, but a 
 correction factor is necessary owing to the following indicated reaction 
 
 Cu SO 4 + Zn Cu + ZnSO 4 . 
 
 OUTLINE OF HYDROGEN. 
 I. History of Hydrogen. 
 
 (a) Discovery when and by whom? 
 
 (b) Derivation of name. (Gr. hudor (water), and. geinomai 
 
 (I produce).) 
 
 (c) Historical in connection with (a) and (b). 
 II. Occurrence or Distribution. 
 
 III. Methods of Preparation. 
 IV. Physical Properties. 
 
 Colorless, Sp. Ht. (gas at const, pres.) 3.409. 
 
 Odorless, Sol'ty in Aq., 1.82 vols. in 160 (20 C.). 
 
 Tasteless, Diffuses rapidly. 
 
 Wt. of il., .09 grm. Melting point, 256 to 257 C. 
 
 Density (O = 1 6), 1.008. Boiling point, 252 to 253 C. 
 
 Density (Air = i ), .0696. Crit. Temp., 238 to 240 C. 
 
 Note. K. Olszewski, in an attempt to liquefy helium, cooled the 
 gas to 259 C. under 180 atmospheres' pressure, by the aid of solid hy- 
 drogen; the pressure was suddenly reduced to that of the atmosphere 
 which should give a degree of cold as calculated by Laplace and Poisson's 
 formula, equal to 271.3 C. Helium did not liquefy. This is the 
 lowest temperature recorded to date. 
 
 V. Chemical Properties. Hydrogen, in the majority of its chemical 
 relationships, displays the characteristics of a metal, and because of this 
 it is usually regarded as a gaseous metal* at ordinary temperatures. 
 
 * NOTE. Since hydrogen does not form a base with oxygen and because its combi- 
 nations with many of the non-metals are acids and not salts, it does not come within 
 the ordinary definition of a metal. 
 
HYDROGEN. 73 
 
 It is practically diametrically opposed to oxygen. It is liberated at the 
 negative electrode by the electric current during electrolysis; it is dis- 
 placed from many of its compounds by metals, i.e., hydrogen functions 
 as a metal; sodium and potassium absorb hydrogen when heated from 
 250 to 400 C., forming alloys (Na^H and K 2 H Richter); similarly the 
 compound PdH 2 conducts itself like an alloy of two metals; according 
 to Graham the specific gravity of the condensed hydrogen in these com- 
 pounds is found to be .62, which makes it somewhat heavier than the metal 
 lithium. Hydrogen combines energetically with oxygen, fluorine, chlor- 
 ine and the metal lithium; however, it unites directly with but few of the 
 elementary substances. Its " affinity" for oxygen and chlorine under 
 certain conditions is such that it will displace the elements with which 
 they are united. When oxygen is taken away from a compound by hydro- 
 gen, the latter is said to be oxidized and the compound reduced. 
 
 The thermochemical deportment of hydrogen and oxygen in the forma- 
 tion of the two compounds, H 2 O and H 2 O 2 , is interesting. 
 
 H 2 + O H 2 O (18 grm. at 20 C.) + 68,360 cals. 
 (H 2 -f O) 3 vols. > H 2 O (2 vols. steam) + 193 cals. 
 H 2 O (steam, 100 C.) H 2 O (water, 100 C.) + 9,666 cals. 
 H 2 O (water, 100 C.) H 2 O (water, 20 C.) + 1,440 cals. 
 193 cals. + 9,666 cals. + 1,440 cals. = 11,299 ca l s - 
 68,360 cals. 11,299 ca l s - = 57>6i cals. 
 
 The approximate thermal equivalent of the chemical energy of a mix- 
 ture of 2 grm. of hydrogen and 16 grm. of oxygen is 57,061 cals. 
 
 The "Law of Hess" is suggested by the thermo-chemical conduct of 
 hydrogen dioxide: 
 
 (i) H 2 + 2 H 2 2 Aq + 45,300 cal. 
 and (2) H 2 O 2 H 2 O + O + 23,100 cal. 
 
 adding (3) H 2 + O H 2 O + 68,400 cal. 
 
 No matter how many stages there are to a given reaction, if the initial 
 and final states are the- same in each case, the heat of the reaction will 
 be constant. 
 
 On inspecting equation (2) it is obvious that H 2 O 2 is an active oxidizer, 
 as it liberates 23,100 cal. more than when the same amount of free oxy- 
 gen is used under similar conditions. This serves to explain its activity 
 as an oxidizer in foregoing experiments. Which is the more active oxid- 
 izer O 3 or H 2 O 2 ? Hint compare energy equations. 
 
 Compounds of Hydrogen. To be studied as work progresses. 
 
 Uses of Hydrogen. : Oxy-hydrogen blowpipe, calcium light, etc. 
 
 Miscellaneous Topics. Principles, theories, definitions, etc. 
 
 Note. In the future the student will make a brief written " resume" 
 of the study of each element. It should embrace the various heads, under 
 which data is to be tabulated, as suggested above in the " Outline of 
 Hydrogen." 
 
74 EXPERIMENTAL CHEMISTRY. 
 
 PROBLEMS. 
 
 i. What is the weight of 100 cm. 3 of hydrogen at 20 C. and 
 760 mm. ? 
 
 2. How many grams of zinc will be required to liberate 10 grams 
 of hydrogen from sulphuric acid? 
 
 3. What is the weight of a liter of hydrogen measured over water at 
 20 C. and 777.36 mm.? 
 
 4. How many liters of hydrogen can be obtained from 100 cm. 3 of 
 sulphuric acid (density, 1.84)? 
 
 5. How much zinc and sulphuric acid will be required (theoretically) 
 to liberate i liter of hydrogen? 
 
 6. A liter of oxygen weighs as much as what number of liters of 
 hydrogen ? 
 
 7. Two grams of hydrogen are equivalent to how many liters of the 
 gas at standard conditions? 
 
 8. What are the valences of the elements in the following: HC1, 
 H 2 O, NasO, NH 3 , KC1, LiH, H 2 S, CO 2 , FeO, Fe^? 
 
 9. If a liter of hydrogen weighs .09 grm. what is the weight of a liter 
 OfO 2 ? OfCO 2 ? OfN 2 ? OfCLj? 
 
 10. What are the relative rates of effusion of hydrogen, oxygen and 
 carbon dioxide? 
 
 ii. Why is the oxygen admitted through the inner tube of the oxy- 
 hydrogen blowpipe? Note: The temperature of the hydrogen flame 
 in air is about 2000 C.; in oxygen it is about 2500 C. It can not sur- 
 pass this latter temperature as it is the temperature at which steam is re- 
 solved into its elements. 
 
 12. What are the formulae of cuprous and cupric oxides? Of fer- 
 rous and ferric oxides? Of ferrous and ferric chlorides? 
 
 13. Is the process of oxidation accompanied by the process of reduc- 
 tion ? Illustrate by equations and interpret. 
 
 14. Is occlusion a chemical or physical action ? State your reasons. 
 
 15. Mention a reaction in which the speed was altered by surface 
 effects. 
 
CHAPTER X. 
 WATER. 
 
 In all investigations, the product of the interaction of two volumes of 
 hydrogen and one volume of oxygen, has been proven to be identical 
 with the substance which is known as water. It is one of the most 
 abundant and universally distributed of all chemical compounds. It is 
 essential to life. Its properties are remarkable and diversified. 
 
 In 1781, Cavendish confirmed the formation of water by the com- 
 bustion of hydrogen. Prior to his work, water was thought to be an ele- 
 mentary substance. Lavoisier determined its quantitative composition 
 in 1783. In 1805, Gay-Lussac showed that it was produced by the union 
 of two volumes of hydrogen with one volume of oxygen. Water was first 
 decomposed by electricity in 1800 by Nicholson and Carlisle. Davy con- 
 firmed and extended the work of these men, aside from initiating many 
 brilliant experiments. The names of the men who were connected with 
 the study of water are, Cavendish, Priestly, Lavoisier, Humbolt, Gay- 
 Lussac, Nicholson and Carlisle, Berzelius, Davy, Dumas, Dulong, Stas, 
 and Morley. 
 
 Experiment I. Composition of Water. 
 
 (a) Synthesis of water. Recall or repeat the various experiments in- 
 volving the " synthesis of water." 
 
 2H 2 (two vols.) + O 2 (one vol.) * 2H 2 O (two vols.) 
 Recall " Law of Combining Volumes." State it. 
 
 (b) Analysis of water. 
 
 1. By electrolysis. 
 
 2. By action of metals. 
 
 4 H 2 (steam) + 3 Fe (hot) <^Fe 3 O 4 + 4 H 2 (Fig. 18). 
 
 At a high temperature, steam oxidizes those elements which readily 
 combine with oxygen. Recall the respective oxidizing activities of O 2 , 
 O 3 and H 2 O 2 . Compare their respective energy equations with that of 
 steam when each is acting as an oxidizing agent. 
 
 Experiment II. Physical Properties. 
 
 (a) Quant.) Density. Clean and dry a 50 cm. 3 Erlenmeyer flask; 
 weigh it. By means of -a burette or a pipette introduce into the flask 15 
 cm. 3 of distilled H 2 O which has a temperature of about 20 C.; weigh 
 flask and contents as rapidly as possible to prevent loss by evaporation. 
 Calculate the approximate density of H 2 O under the existing conditions. 
 See "Table of Density" in Appendix. 
 
 75 
 
70 EXPERIMENTAL CHEMISTRY. 
 
 Note. Burettes filled with distilled H 2 O may be placed in readily 
 accessible places in the laboratory. 
 
 (b) Freezing point. Fill a beaker with a mixture of clean ice and water; 
 the ice should be broken into small pieces. Suspend a thermometer in 
 the mixture for 5 or 10 min.; the mixture should be stirred. Tap the 
 thermometer with finger and read. Record reading. 
 
 (c) Boiling point. Measure 20-25 cm - 3 f distilled H 2 O into a clean 
 Erlenmeyer flask; suspend a thermometer in the flask so that its bulb is 
 about 2 or 3 cm. above surface of H 2 O; heat flask and contents until there 
 
 FIG. 18. (Smith and Keller.) 
 
 is a rapid evolution of steam; tap thermometer, and read. Record 
 reading. Suspend thermometer so that bulb dips into water but does 
 not touch bottom of flask. Repeat above. Record reading. Read 
 the barometer. (Instructions.) Record reading. 
 
 Experiment III. Purity of Water. 
 
 Place a few drops of distilled water upon a clean "watch glass" and 
 evaporate to dryness upon a steam bath. Is there a stain or residue on 
 glass ? Repeat above using ordinary water, for example, " drinking water." 
 Compare results with above. If the impurities of the water are very vola- 
 tile will this method enable you to detect them ? Why ? 
 
 The preparation of "absolutely" pure water is an impossibility, as the 
 material of any vessel is soluble in a greater or lesser degree. 
 
 Experiment IV. "Temperature -Density " Graph of Water. 
 
 By referring to the table in the Appendix, plot on "coordinate paper" 
 the "temperature-density" graph of water. Use the axis of ordinates for 
 the scale of density and the axis of abscissas for temperature-scale. 
 (Instructions.) 
 
 Does the comparatively abrupt change in the direction of graph sug- 
 gest the introduction of a new "factor" in the phenomena which graph is 
 supposed to represent? May this change in direction merely represent 
 that a factor or a group of factors has suddenly become dominant? 
 
WATER. 77 
 
 Experiment V. Purification of Water. 
 I. Chemical Methods. 
 II. Mechanical Methods. 
 Decanting. 
 Filtering. 
 Boiling. 
 Distillation. 
 
 (a) Distillation of a solution of copper sulphate. Assemble a distilling 
 apparatus, composed of a distilling flask, Liebig condenser (Fig. 19) and 
 a receiver, or use a retort and a receiving flask. (Instructions.) Test 
 solution with litmus paper; test distillate with litmus paper; compare 
 results. How does color of distillate compare with color of original 
 solution ? 
 
 FIG. 19. 
 
 (b) Distillation of a solution of ammonium hydroxide. Proceed as in 
 (a) Results? Explain. 
 
 (c) Distillation of a solution of alcohol. Proceed as in (a), but test in- 
 flammability of solution and distillate by means of a lighted match. 
 Explain. Recall "Law of Partial Pressures." 
 
 Note. Boiling-tubes prevent "bumping" of solutions undergoing 
 ebullition. 
 
 Experiment VI. Hydrolysis. 
 
 To a small quantity of bismuth chloride, Bi C1 3 , or antimony chloride, 
 SbCl 3 , add 5 cm. 3 of water; the white precipitate is essentially bismuth 
 oxychloride, BiOCl. To the precipitate add a few drops of concen- 
 trated hydrochloric acid, then warm tube; by repeating this process 
 
78 EXPERIMENTAL CHEMISTRY. 
 
 dissolve the precipitate in a minimum quantity of acid. The tube con- 
 tains a solution of SbCl 3 . To a test tube nearly full of water add a 
 few dr,ops of the solution. Explain the formation of the white precipitate. 
 Write equations for the three reactions. 
 
 Experiment VII. Chemical Union of Water with Oxides. 
 
 (a) Combination with a metallic oxide. Place a small quantity of 
 quick lime (CaO) or barium oxide (BaO) in a test tube; add 5-10 cm. 3 
 of H 2 O which has been previously tested with litmus paper; shake vigor- 
 ously; let contents settle; decant clear liquid into another test tube and 
 test liquid with litmus paper. Results ? Equation ? Substances like 
 the above which turn red litmus blue are called "bases.'' 
 
 (b) Combination with a non-metallic oxide. Recall the action of P 2 O 5 
 and SO 2 on the moist litmus paper. Record results. Substances formed 
 by the union of non-metallic oxides and water and which turn blue 
 litmus red, are called "acids." 
 
 Experiment VIII. Water of Hydration (Crystallization). 
 
 (a) Place a small crystal of copper sulphate (CuSO 4 .5H 2 O) in a test 
 tube; heat gently until a white powder remains. Is there evidence that 
 water has been liberated during the above process ? Add a drop or two of 
 water to the powder when tube is cool. Effect ? Add water and boil, dis- 
 solving the powder in the least possible quantity of water; set tube aside 
 for several days; do crystals form? 
 
 (b) Using crystals of gypsum, potassium dichromate, barium chloride 
 and potassium nitrate, ascertain whether water is present in each crystal. 
 
 Note. Some crystals contain " mechanically inclosed" water; when 
 such crystals are heated they fly to pieces explosively: they are said to 
 decrepitate. Did any of the above crystals decrepitate? 
 
 Do all crystals contain water oj hydration ? Reasons for your answer ? 
 
 A salt containing water of hydration is spoken of as a hydrated salt; 
 when the water has been removed it is known as an anhydrous or dehy- 
 drated salt. 
 
 Experiment IX. Vapor Tension of Substances. Efflorescence and 
 Deliquescence. 
 
 (a) Examine a small clear crystal of sodium sulphate, Na.jSO 4 .io H 2 O; 
 place it on a clean watch glass and set it in closet; after several days ex- 
 amine. Explain. Equation ? The substance is said to be efflorescent. 
 
 (b) Place a piece of dehydrated calcium chloride, CaCl 2 in a beaker. 
 Repeat (a). Explain. Equations? The substance is said to be 
 deliquescent. 
 
 (c) Introduce 10 cm. 3 of concentrated sulphuric acid into a dry test 
 tube; mark its height by means of a piece of label or an ink mark; allow 
 it to stand for a week. Explain. 
 
WATER. 79 
 
 Experiment X. (Quant.) Determination of Water of Crystallization. 
 
 Weigh accurately a clean dry crucible and cover. Record weight. 
 Introduce into crucible about 2 grm. of powdered copper sulphate crys- 
 tals ; place cover on crucible and weigh accurately. Record weight. Place 
 covered crucible on a pipe-stem triangle and heat gently for about 25 min. 
 or until quite certain that the blue color has completely disappeared; 
 cool in desiccator, then weigh. Record weight. Heat again for 5-10 
 min. and weigh. Repeat until weight becomes constant. Calculate 
 the percentage of water of crystallization of the crystals. 
 
 Note. In applying heat to the crucible, the tip of the flame may 
 barely touch the bottom of the crucible. 
 
 Physical Properties of Water. 
 
 Tasteless. Ht. of Fusion, 80 cals. 
 
 Odorless. Ht. of Vaporization, 536.7 cals. 
 
 Colorless (in thin layers). Melting point, o C. 760 mm. 
 Density (i cm. 3 at 4 C.), i grm. Boiling point, 100 C. 760 mm. 
 
 Sp. Ht. (Solid state), .50. Crit. temp., 370 C. (Highest critical 
 Sp. Ht. (Liquid), i.oo temperature known.) 
 
 Sp. Ht. (Gas), .477- 
 
 The specific heat of water is remarkably high. More heat is required 
 to raise the temperature of a given weight of it one degree than is re- 
 quired for any other substance except hydrogen. 
 
 Chemical Properties. Water is one of the most stable of all substances, 
 i.e., it is not easily decomposed as regarded from the standpoint of energy, 
 yet like other chemical compounds it is broken up into its elements by 
 heat. "Sainte- Claire Deville was the first to carefully investigate and 
 explain the decomposition of water by pouring molten platinum (1770 C.) 
 into it." " He proved that the dissociation a reversible decomposition 
 did not take place suddenly, but gradually; that it advanced regularly 
 with increasing temperature, and was limited by an opposing combina- 
 tion-tendency on the part of the components." Richter. 
 
 The decomposition is appreciably initiated at 1000 C. and is about half 
 complete at 2500 C. The percentage of dissociation increases with in- 
 crease of temperature. 
 
 Water bears evidence of being slightly dissociated when in the liquid 
 condition as well as when in the gaseous state. Electrical conductivity 
 and certain chemical reactions indicate the truth of the following equation: 
 
 H 2 O<=H+ + OH'. 
 
 This will be discussed more fully under the subject of "The Modern 
 Theory of Solution." Water is usually referred to as a perfectly neutral 
 substance, but this is far from agreeing with experimental facts. Owing 
 
8o EXPERIMENTAL CHEMISTRY. 
 
 largely to the products of dissociation, water manifests definite chemical 
 properties under definite conditions. 
 
 Another interesting property of water is one which is thought to be 
 closely related to the varying density of water with change of temperature. 
 In virtue of this property the molecules of water are supposed to associate, 
 that is, a number of molecules combine and form a " molecular-complex." 
 This phenomenon is known as the polymerization of water. " Surface 
 tension" and "depression of the freezing point" experiments support this 
 view. Raoult and others hold that at temperatures near the freezing 
 point, a relatively large quantity of the water possesses the molecular 
 formula, H 8 O 4 or (H 2 O) 4 . Increase in temperature causes a breaking 
 up of these molecular-complexes, and vice versa. The change in density 
 of water with change of temperature is explained in general as follows: 
 
 The H 8 O 4 molecules occupy more space than 4H 2 O; polymerization 
 increases as temperature is lowered; water contracts when the temperature 
 falls; two factors must then be considered. Water expands when cooled 
 below 4 C. because the effects of polymerization overbalance the effects 
 due to contraction as the result of lowering the temperature. When the 
 temperature is increased to 4 C., the water should expand normally, 
 and probably does, but the molecular-complexes tend to split up which 
 per se causes a shrinkage in volume; at 4 C. the effects of the two factors 
 balance each other, hence the maximum density at this temperature. If 
 the temperature is raised above 4 C., then the expansion-effects due to in- 
 crease of temperature are greater than the contraction-effects due to a 
 breaking down of the associated molecules. Above 4 C. the effect of 
 cubical expansion due to increased temperature is dominant; below 4 C., 
 expansion-effects due to increased polymerization are dominant. 
 
 It was said in the beginning of this chapter that water in its chemical 
 relations presents some very remarkable features; one of the most marked 
 is, that although it is an indifferent oxide (hydrogen oxide), it possesses 
 a group of combining tendencies which extend over a wider range than 
 those of any other chemical compound. It combines directly with many 
 compound substances, but with few elements chlorine and bromine being 
 perhaps the only elementary substances. Gases are dissolved by water 
 to a relatively slight extent, but it can not be affirmed positively that they 
 enter into a chemical combination. The most common of its chemical 
 reactions is the formation of a class of compounds known as hydrates, 
 which exist in the solid form only, undergoing decomposition when placed 
 in solution, Ex.s. CuSO 4 .5H 2 O, Na^SO^io H 2 O. They show a definite 
 composition and frequently much heat is developed during their for- 
 mation, thus (Na^COg, 10 H 2 O), equals 8800 cal. Water also reacts with 
 some substances in a manner which is very similar to metathesis to say 
 that the products of dissociated water, rather than the water per se, inter- 
 acts with substances to produce double decomposition, is a statement more 
 nearly in accord with facts. This kind of a reaction is known as hydrol- 
 ysis. The mechanism of the interaction may be interpreted more readily 
 by the aid of an equation. 
 
WATER. 8 1 
 
 H 2 <= H+ + OH 
 
 3d' + 2H+ + 2QH / <=>Bi(OH 2 )Cl + 2 H + ' + 2 Cl' 
 Bi(OH) 2 Cl<=BiOCL + H 2 O. 
 
 If the product of hydrolysis is soluble there is usually little or no visible 
 evidence of the interaction; if the product is insoluble a precipitate forms. 
 As the work advances the phenomenon of hydrolysis will be more thor- 
 ough discussed. A familiar example of the direct combination of water 
 and metallic and non-metallic oxides is the slaking of quicklime which 
 may be represented as follows: 
 
 CaO+ H 2 O Ca(OH) 2 (Calcium Hydroxide), 
 again, SO 2 + H 2 O H 2 SO 3 (Sulphurous Acid). 
 
 The product of the first reaction belongs to a group of substances known 
 as bases; the latter product, to a group whose generic name is acids. 
 The chemical properties of these two groups, acids and bases, are very 
 different; an aqueous solution of the former turns red litmus paper blue, 
 while in the case of the acids, blue litmus paper is turned to a red color. 
 
 However, of all the various properties of water, perhaps none are of 
 more importance to the chemist than its solvent properties. The question 
 as to whether the effecting of a solution is a chemical or physical process 
 has not been satisfactorily answered. This question will be discussed 
 more fully in the next chapter. 
 
 In addition to its enumerated uses, it should be remembered that water 
 is the standard of many physical measurements . 
 
CHAPTER XI. 
 SOLUTIONS. 
 
 We have observed that many gases, liquids and solids when placed 
 in water disappear and form homogeneou systems which are known as 
 solutions. The operation of preparing a solution is called "dissolving" 
 or " putting into solution." The substance dissolved is known as the 
 solute and the material in which the solute is dissolved, the solvent. 
 These two terms are unfortunate inasmuch as they do not suggest the 
 mutual interaction of solute and solvent during the process of dissolving. 
 All substances are soluble in a degree, the solubility depending upon the 
 relative strength of the affinities oj the substance for itself and for the solvent. 
 If the affinity of solvent and solute is greater than the affinity of sub- 
 stance for self, then the solubility will be correspondingly great. The 
 degree of solubility is usually expressed by the terms insoluble, slightly 
 soluble, soluble, and very soluble. When the solution contains a rela- 
 tively large quantity of the solute it is said to be a concentrated solution; 
 if a relatively small quantity, a dilute solution. 
 
 " Since matter in every state can be mixed with other matter, irre- 
 spective of its state, it is obvious that many different kinds of solutions 
 are possible." Jones (H. G.) gives the following list: 
 
 I. Solution of a solid in a solid. 
 II. Solution of a solid in a liquid. 
 
 III. Solution of a solid in a gas. 
 
 IV. Solution of a liquid in a solid. 
 V. Solution of a liquid in a liquid. 
 
 VI. Solution of a liquid in a gas. 
 
 VII. Solution of a gas in a solid. 
 
 VIII. Solution of a gas in a liquid. 
 
 IX. Solution of a gas in a gas. 
 
 Among the properties of solutions may be noted color, odor, taste, 
 density, expansibility, compressibility, surface tension, viscosity, vapor 
 tension, osmotic pressure, refractive index, definite boiling and freezing 
 points, etc. 
 
 SOLUBILITY OF SOLIDS. 
 
 Experiment I. Characteristics of a Solution. 
 
 Powder separately a few small crystals of alum, and potassium di- 
 chromate; dissolve each in the least possible quantity of water the test 
 tubes should be shaken and warmed repeatedly to aid in the process. 
 Are the solutions clear? Transparent? Homogeneous? Pour the 
 solutions into separate crystallizing dishes, or set the tubes aside to cool. 
 
 82 
 
SOLUTIONS. 83 
 
 Do crystals form ? If so, decant the liquid and examine crystals. Are 
 they similar to the original crystals ? 
 
 A solution is ordinarily defined as a clear transparent homogeneous 
 mixture, the components of which can not be separated by a purely 
 mechanical process. "Colloidal" or "pseudo-solutions" see Ostwald's 
 "Principles of Chemistry." to 
 
 Experiment II. Surface and Diffusion Phenomena. Solution Tension. 
 
 Half fill a test tube with water; drop one small crystal of potassium 
 permanganate into the water; set tube where it will not be shaken, yet 
 can be easily observed; as you continue your experimenting, notice the 
 color of the solution at short intervals of time. Record observations. 
 Was a relatively long period of time required for the dissolving of the 
 crystals ? Postulating that the salt is quite soluble in water at a given 
 temperature, what factors determine largely the speed of the dissolv- 
 ing process? To what two mechanical processes do we resort 
 usually in an endeavor to hasten solution? Did any phenomena occur 
 during the dissolving of the solid which would suggest any of the prop- 
 erties possessed by liquids or gases ? Enumerate them. 
 
 The process of solution receives a partial explanation, at least, when 
 the accompanying phenomena are interpreted in terms of the kinetic- 
 molecular hypothesis. When a soluble crystalline substance is intro- 
 duced into the solvent, it is thought that the molecules of which the 
 solute is composed are detached and enter the solvent. By the process 
 of diffusion these detached particles move away from the surface of the 
 solute, and are scattered throughout the solvent. After a time some of 
 these particles, which move in every direction, will again come into con- 
 tact with the solid solute and attach themselves to it. If the rate at which 
 the molecules press into solution is greater than that with which they 
 return to the solute, it is evident that the solute will eventually be wholly 
 dissolved. If a sufficient quantity of the solid substance has been placed 
 in the solvent the dissolving process will continue until the speeds of the 
 two opposing actions are identical. When this occurs the solution is 
 said to be saturated. The solid solute is then in equilibrium with the 
 dissolved portion, 
 
 KMn0 4 (Solid) <=> KMnO 4 (Dissolved). 
 
 If this equilibrium is disturbed by varying the concentration of the 
 solution, the speeds of the opposing actions will be altered until equi- 
 librium is again established, either by the dissolving of more of the 
 solute if the concentration is decreased, say by diluting the solution; 
 or by the deposition of a portion of the solute if the concentration is 
 increased by any cause whatever say the removal of a portion of the 
 solvent. The experimenter is never assured of the existence of an 
 equilibrium unless a portion of the solute (solid) is in contact with the 
 
84 EXPERIMENTAL CHEMISTRY. 
 
 solution. Powdering the solute and shaking the mixture hastens the 
 process of solution. Why? 
 
 This tendency of the molecules of a substance when placed in a solvent, 
 to leave the solid and pass into solution is referred to as solution tension 
 because of its evident analogy to the vapor tension of liquids, in virtue 
 of which liquids tend to assume the gaseous condition. The analogy is 
 continued: There is equilibrium when the vapor tension of a liquid is 
 balanced by the gaseous pressure of the vapor above it; likewise there is 
 a dynamic equilibrium in a saturated (concentrated) solution between 
 the dissolved and undissolved portions of the solute; the force (energy) 
 in virtue of which the molecules tend to pass into solution is usually 
 spoken of (mentioned previously) as the solution tension of the solute; the 
 force in equilibrium with the solution tension, and which is considered 
 as the analogue of the gaseous pressure of the vapor above a liquid, is 
 known as osmotic pressure. In this pressure we recognize the cause of 
 the diffusion of substances in solution. It has been discovered that the 
 molecules of a dissolved substance like sugar, exert a pressure on the 
 solvent identical with the pressure which they would exert on the sides 
 of a vessel of the same volume as that of a solution, if they were in the 
 gaseous state. This pressure is given the name osmotic pressure because 
 it is only by taking advantage of the phenomenon of osmosis (which 
 provides for the elimination of "surface pressures") that it can be 
 rendered apparent and directly measured. 
 
 The correlation of much experimental data which will be discussed 
 as the work proceeds, has given rise to what is known as the " physical 
 theory" of dilute solutions. It may be stated as follows: "The mole- 
 cules of the dissolved substance pervade the solvent without being 
 influenced thereby, and possess the same properties as they would 
 possess did they alone, in the state of gas occupy the volume filled by 
 the solution." Bloxam. It seems that there is more likelihood of the 
 dissolved substance being in a condition comparable to a gas than to 
 either a liquid or a solid. It is obvious that it can not be in a state of 
 aggregation comparable to a solid. 
 
 Hulett (J. Amer. Chem. Soc., 27, 49, 1905) recently presented an 
 interesting article on the importance of the state of the solid from which 
 the solution is made. The real subject of the article, however, is an 
 explanation of the greater speed of solution and greater solubility of very 
 small particles of a solid as "based on the following considerations: 
 The boundary between a solid and a liquid is the seat of a certain amount 
 of energy due to the surface-energy of the liquid; if this surface is increased 
 by powdering the solid, the total surface-energy is correspondingly in- 
 creased. Further, it is a generally observed fact that the form of a sub- 
 stance which has the greater free energy is the more soluble, has the 
 greater vapor-pressure, and is the least stable form, e.g., allotropic 
 modifications of substances have different solubilities, and the unstable 
 form is always the more soluble. This phenomenon js hardly analogous 
 to the well known behavior of liquid drops of different sizes. Small 
 
SOLUTIONS. 85 
 
 drops in the vicinity of large ones grow small and disappear, while the 
 larger ones grow larger, and the reason is quite clear. It is known that 
 the curved surface of a liquid has a greater vapor-pressure than a plane 
 or less curved surface; therefore, a distillation takes place. The simi- 
 larity between the vapor-pressure of liquids and the solution-pressure 
 of solids has suggested to some the analogy between the facts just men- 
 tioned and the behavior of solid particles of different sizes in contact 
 with the solution. But we can not assume that the surface of the par- 
 ticles of a powder is curved, or, if that is granted, we do not know that a 
 curved surface of a solid or a sharp edge has a greater solution-pressure 
 than a plane surface of the same substance. 
 
 "Hulett found that a solution of gypsum saturated at 25, containing 
 2.080 grm. CaSO 4 per liter will increase its concentration rapidly to a 
 maximum when shaken with powdered gypsum and then will decrease 
 to the original value again. In one experiment the content of gypsum 
 reached 2.542 grm. CaSO 4 in a liter in a minute. The fine powder used 
 for. this purpose was found, after the concentration had reached its 
 original point, to have increased in size of its grain. The smallest 
 particles thus go into solution, produce the supersaturation and are 
 then deposited upon the others, so that all increase in size." Morgan. 
 Hulett suggests the preparation of saturated solutions from large particles 
 as there is less likelihood of supersaturation. 
 
 OSMOSIS. 
 
 " If a solution and the pure solvent are separated by a semipermeable 
 membrane the solvent will flow through the membrane into the solution, 
 where its escaping tendency is less. The only way of preventing this 
 flow is to make the escaping tendency of the solvent the same on both 
 sides of the membrane. There are two simple ways of accomplishing 
 this, (i) to increase the pressure on the solution until the escaping 
 tendency of the solvent in the solution is raised to equal that of the 
 solvent in the pure state, (2) to diminish the pressure on the pure solvent 
 until its escaping tendency is lowered to equal that of the solvent in the 
 solution. 
 
 The osmotic pressure may, therefore, be defined in two ways, (i) as 
 usually defined, it is the increase in the pressure on the solution necessary 
 to bring the latter into equilibrium with the solvent; (2) Noyes,* however, 
 prefers to define the osmotic pressure as the diminution in the pressure 
 on the solvent necessary to bring it into equilibrium with the solution. 
 Neither of these definitions is entirely free from objections." Lewis, G. N.f 
 
 Experiment III. (L. T.) Osmotic Pressure. 
 
 Note. The instructor or assistant should assemble the apparatus 
 (Fig. 20) in the presence of the members of the class. The use of a solu- 
 
 *Z. physik. Chem., 35, 707 (1900). 
 
 t Osmotic Pressure of Concentrated Solutions." Jour. Amer. Ch3m. Soc., 30, 668 
 (1908). 
 
86 
 
 EXPERIMENTAL CHEMISTRY. 
 
 tion of sugar is suggested. Whether the experiment is designed to show 
 relationships in a quantitative or qualitative manner, the effects of concen- 
 tration and temperature upon the osmotic pressure should be demon- 
 strated. These effects should be compared with corresponding effects 
 
 FIG. 20. 
 
 produced on gases by similar influences. By use of the following data 
 the possibility of applying the "gas law," PV = RT, to dilute solutions, 
 may be shown. 
 
 OSMOTIC PRESSURE OF CANE SUGAR. 
 
 (" Osmotic Investigations "- 
 
 Ejject oj Concentration. 
 
 Concentration in Pressure in 
 
 per cent, by weight. mm. of Hg. 
 
 1 per cent. 535 mm. 
 
 2 per cent. 1016 mm. 
 2.74 per cent. 1518 mm. 
 4 per cent. 2052 mm. 
 6 per cent. 375 mm. 
 i per cent. 535 mm. 
 
 -Pfeffer, Ames Sci. Mem's.) 
 
 Effect of Temperature. 
 (i% sol. of cane sugar.) 
 Temperature. Pressure. 
 14.2 C. 510 mm. 
 
 32.0 C. 544 mm. 
 
 6.8 C. 
 I 3-7 C. 
 22.0 C. 
 
 505 mm. 
 525 mm. 
 548 mm. 
 
SOLUTIONS. 
 
 It will be seen that Pfeffer found the osmotic pressure of a i per cent, 
 sugar solution at 6.8 C. to be equal to 50.5 cm. of mercury or 50.5 x 
 13.59 grm. per sq. cm. A i per cent, sugar solution contains approx- 
 imately i grm. of sugar per 100 cm.s of solution. As the molecular 
 weight of sugar (C 12 H 22 O U ) is 342, then a gram-molecular weight is 
 contained in 34,200 cm. 3 of the solution. This volume represents then 
 the molar volume at 6.8 C. or 279.8 C. on absolute scale (T). 
 
 PV = RT, 
 
 PV 
 
 or, - = R = 84800 gr. cms. 
 
 T 
 
 50.5 x 13.59 x 34200 
 
 Substituting, - = 83,900 (approx.). 
 
 279.8 
 
 OSMOTIC PRESSURE AND MOLECULAR WEIGHTS. 
 
 (Sugar Solutions at about 20 C.) 
 (Morse and Frazer.) 
 
 Weight 
 
 Volume 
 
 normal 
 
 normal 
 
 Moles in 1000 
 
 Moles per 
 
 grm. H 2 O. 
 
 Liter. 
 
 (W) 
 
 
 0.05 
 
 0.04948 
 
 o. 10 
 
 0.09794 
 
 o. 20 
 
 o. 19192 
 
 0.25 
 
 0.23748 
 
 0.30 
 
 0.28213 
 
 0.40 
 
 0.36886 
 
 0.50 
 
 0.45228 
 
 0.60 
 
 0.53252 
 
 0.70 
 
 0.60981 
 
 0.80 
 
 0.68428 
 
 0.89101 
 
 0.75000 
 
 0.90 
 
 0.75610 
 
 1 .00 
 
 0.82534 
 
 Pressures at Same 
 
 Temperature. W(22.4 
 
 M=- 
 Gaseous. Osmotic (P). 
 
 o.o824t). 
 
 i . 21 
 
 2.40 
 
 4.82 
 
 6.06 
 
 7.22 
 
 9.68 
 
 12.07 
 
 14-58 
 
 17. 16 
 
 19.17 
 
 21.48 
 
 21.73 
 
 24.27 
 
 1.26 
 2.44 
 4-78 
 6-05 
 
 7-23 
 9.66 
 12.09 
 14-38 
 17.03 
 19-38 
 
 21 . 21 
 21. 8l 
 24.49 
 
 Mean, 341.2 
 
 Van't Hoff summed up these results in the form of a law which bears 
 his name: " The osmotic pressure of a substance in solution is the same 
 pressure which that substance would exert were it in gaseous form at the 
 same temperature and occupying the same volume." 
 
 In a paper which appeared in the Amer. Chem. Jour., in July, 1905, 
 Morse and Frazer show as the result of a series of accurate experiments, 
 that Van't Hoff's law holds for solutions of sugar if the words, "volume 
 
88 EXPERIMENTAL CHEMISTRY. 
 
 of the pure solvent," are substituted for the word "volume," which refers 
 to the total volume of the solution. They say: "When we dissolved 
 a gram-molecular weight of cane sugar (342.22 grm.) in 1000 grams of 
 water, i.e., in that mass of solvent which has the unit volume, i liter, at 
 the temperature of maximum density, we found its osmotic pressure, 
 at about 20, in quite close accord with the pressure which a gram- 
 molecular weight of hydrogen would exert, at the same temperature, 
 if its volume were reduced to i liter, i.e., to that volume which the unit 
 mass of solvent has at the temperature of greatest density." Or in 
 other words, a substance in solution " exerts an osmotic pressure through- 
 out the larger volume of the solution equal to that which as a gas it would 
 exert if confined to the smaller volume of the pure solvent." 
 
 " It should be borne in mind," says Walker, "that the osmotic pressure 
 in a solution may be regarded as always present, whether a semiper- 
 meable membrane renders it visible or not. The osmotic pressure in 
 the ordinary reagent bottles of the laboratory is of the dimensions of 
 50 atmospheres. This pressure is, of course, not borne by the walls 
 of the bottle nor is it apparent at the free surface of the liquid. Where 
 the liquid comes in contact with the enclosing vessel there we find a 
 liquid surface, and a consideration of the magnitude of the forces at work 
 in the phenomena of surface tension leads us to believe that the pressure 
 at right angles to the free surface of a liquid, and directed towards the 
 interior of the liquid, is measurable in hundreds and even thousands of 
 atmospheres. Osmotic pressures, then, large as they are in ordinary 
 solutions, are small compared to the surface pressures in liquids, and 
 their existence is consequently not evident at the free surface of liquids. 
 It is only when these surface pressures are got rid of that we can measure 
 osmotic pressures directly. The liquid solvent can easily penetrate the 
 semipermeable membrane, so at the semipermeable membrane there is 
 no surface pressure of the ordinary type. This continuity of the liquid 
 through the semipermeable partition gives us, therefore, the opportunity 
 of determining differences of internal pressure in the solution and the 
 solvent. Various hypotheses have been put forward to explain the nature 
 of osmotic pressure, but none of them can be accounted satisfactory." 
 
 Experiment IV. Effect of Temperature on Solubility of Solids. 
 
 (a) Temperature increases solubility. Add sodium chloride to a 
 test tube half filled with water; shake and continue to add salt until no 
 more will dissolve. The solution is said to be saturated at the temper- 
 ature of the solution. Heat the test tube and contents; add salt until 
 solution is saturated at the higher temperature; set tube aside to cool 
 if there is no undissolved salt in it, otherwise, filter and observe filtrate 
 on cooling. Is the supernatant liquid saturated at the prevailing tem- 
 perature? Decant a portion of the clear liquid into another test tube; 
 cool the contents to a lower temperature by immersing tube in a mix- 
 ture of ice and water, or allow cold tap water to drip upon tube. Is 
 more salt deposited? Is it in equilibrium with the salt in solution? 
 
SOLUTIONS. 8() 
 
 Can the equilibrium be destroyed temporarily by adding water? Try. 
 Explain. Boil solution in test tube until salt is deposited. Explain. 
 
 (b) Temperature diminishes solubility. Prepare a concentrated 
 solution of calcium citrate at the temperature of the laboratory; heat, 
 but do not boil, as it is desired to avoid the vaporization of any appre- 
 ciable amount of water. Results ? Explain . 
 
 (c) Repeat (b) using calcium hydroxide. 
 
 Experiment V. Supersaturated Solutions. 
 
 Fill a test tube of medium size nearly full of crystallized sodium sul- 
 phate; add 4 or 5 cm. 3 of water and heat gently until solution has a tem- 
 perature of about 30 C.; shake; add salt until a saturated solution is 
 procured; pour solution into a clean, dry test tube or small flask; cover 
 the vessel; allow it to cool, then introduce a small crystal of sodium 
 sulphate. If solution has been prepared properly the excess of salt 
 will crystallize out of the solution. Is the supersaturated solution a case 
 of stable or unstable equilibrium when in contact with the solid solute ? 
 
 Experiment VI. Thermal Phenomena Accompanying the Dissolving 
 of a Solute in a Pure Solvent. 
 
 Measure 10 cm. 3 of water into each of five test tubes; take the tem- 
 perature of the water in each tube; to the water in one of the tubes add 
 slowly 5 grm. (see sp. gr., do not weigh) of concentrated H 2 SO 4 , stirring 
 carefully with the thermometer as acid is added. When the contents 
 of tube are homogeneous, i.e., one phase, record the reading of the 
 thermometer; remove thermometer and clean it. 
 
 Repeat above using separately, 5 grm. of solid ammonium chloride 
 (NH 4 C1), dehydrated CuSO 4 , Na^C^ic H 2 O, Na^SO,. Tabulate data. 
 Conclusions ? 
 
 Experiment VII. Relative Solubility of Solids. Effect of Temperature. 
 
 (a) Place i grm. of CuO in a test tube; add 6 cm. 3 of water. Does it 
 dissolve? Heat the mixture to boiling. Effect? 
 
 (b) To 5 grm. of NaCl add 6 cm. 3 of water. Does the salt dissolve? 
 Heat mixture to boiling in an endeavor to dissolve all of the salt. Are 
 you successful? Pour the hot saturated solution upon a dry filter; 
 collect the filtrate in a test tube and cool. Results? Explain. 
 
 (c) Pulverize about 15 grm. of K 2 Cr 2 O 7 ; to 5 grm. in a test tube add 
 6 cm. 3 of water. Will all of the solid dissolve in the solvent ? Heat 
 mixture to boiling. Has all of the solid dissolved? Cool the solution. 
 Observe the effect of lowering the temperature. 
 
 Note. The pulverized K 2 Cr 2 O 7 will be needed in Exp. VIII. 
 
 Experiment. VIII. (Quant.) Prepare a saturated solution of K 2 Cr 2 O 7 ; 
 to 8 grm. of the powdered substance in a flask or beaker, add 50 cm. 3 of 
 distilled water; assist the process of dissolving by frequently shaking it. 
 If the directions have been adhered to the solution will probably be 
 
90 EXPERIMENTAL CHEMISTRY. 
 
 saturated at the end of 10 min. Take the temperature of the solution. 
 Into a weighed evaporating dish, weigh accurately 20-30 grm. of the 
 solution; evaporate to dryness; cool; weigh. Repeat heating and weigh- 
 ing. Calculate the weight of dichromate in i liter of the solution (satur- 
 ated) at the observed temperature. How many gram-molecules (moles) 
 of potassium dichromate in a liter of the saturated solution ? 
 
 Terminology of Solutions. When a substance dissolves in a liquid 
 there is for each temperature and pressure a definite solubility, i.e., a 
 definite relation between the solute and the solvent. If the solution con- 
 tains less of the solute than corresponds to the latter's solubility in the 
 solvent, it is unsaturated; if the amount in solution is in excess of the 
 amount required to saturate it, the solution is said to be supersaturated. 
 The test of the degree of saturation of a solution is made by placing a 
 portion of the solid solute in contact with it; if the solution is unsaturated, 
 a portion of the solute will dissolve; if supersaturated, a portion of the 
 solute will separate from the solution, and continue to be deposited until 
 the solution is saturated and a condition of stable equilibrium is estab- 
 lished. 
 
 Although the above nomenclature is convenient it does not convey 
 the definite information which is so much desired by the chemist. Be- 
 cause of this, the concentrations of solutions are frequently expressed in 
 terms of physical or chemical units. 
 
 The concentration of any substance is the total amount of that sub- 
 stance in solution in a unit volume. 
 
 The solubility of a substance is expressed in terms of the number of 
 grams or gram-molecules which can be dissolved in a unit volume at a 
 given temperature and pressure. 
 
 A solutio'n is referred to as being standard when its concentration is 
 known. Of more frequent use, however, are the terms, normal solutions 
 and molar solutions. 
 
 A normal solution is a standard solution which contains in one liter the 
 hydrogen equivalent of the active reagent, expressed in grams. Thus 
 a normal solution of HC1 contains 36.45 grams of hydrogen chloride; 
 
 1TT0 ^ Mol. Wt. , TT Mol.Wt. . 
 
 normal H 2 SO 4 , - grams; normal KOH, grams; normal 
 
 iodine, 126.97 grams in a liter. If a liter contains T V of an equivalent 
 weight it is designated tenth or deci-normal (.iN); if T J~g- of an equivalent, 
 hundredth or centi-normal (.oiN). 
 
 A molar solution is a standard solution which contains one mole or one 
 gram-molecular weight of the solute in one liter of solution. Thus a 
 molar solution of HC1 contains 36.45 grm. of hydrogen chloride; molar 
 H 2 SO 4 , 98.07 grm.; molar NaOH, 40.05 grm. of sodium hydroxide. 
 
 Influence oj Temperature On Solubility oj Solids. It is a general rule, 
 if the temperature changes, the solubility changes. The solubility of 
 the majority of solid substances increases with increase of temperature; 
 however, there are cases in which the solubility is decreased by rise of 
 
SOLUTIONS. 
 
 temperature. The increase in the solubility of potassium chloride is 
 approximately proportional to the increase of temperature. The solu- 
 bility of calcium citrate is greater at 10 C. than it is at higher temper- 
 atures, say 70 C.; the same is true of calcium hydroxide. It is quite 
 probable that in the majority of cases there is a fall of temperature due 
 to the mere act of solution, but the heat of the chemical combination 
 which undoubtedly follows the process of dissolving in many instances, 
 is frequently so much in excess, that heat alone is the observed result 
 of solution. For example, when sulphuric acid is dissolved in water, 
 it combines with the water to form hydrates of sulphuric acid and a large 
 quantity of heat is developed; the same conduct is displayed by dehy- 
 drated salts, like Na 2 CO 3 and N^SO^ The question, is there a relation 
 between the influence of temperature upon solubility, and some other 
 property of substances, may be answered by an application of that 
 principle which is the basis of explanation of all influences affecting 
 equilibrium; the principle referred to is the one enunciated in the form 
 of La Chatelier's Theorem. If a substance dissolves with an absorp- 
 tion of heat, its solubility will increase with rise of temperature; if, on the 
 other hand, heat is developed on solution, solubility will decrease with 
 lise of temperature. When substances dissolve without thermal alter- 
 ations, solubility is practically independent of temperature. The 
 solubility of sodium chloride is affected but little by temperature alter- 
 ations. A simple test of the influence of temperature upon the solubility 
 of a given substance in a particular solvent may be made by preparing 
 a saturated solution at a given temperature, then observe if increase in 
 temperature causes deposition of the solid; if solute is not deposited, 
 add a small crystal of it to the solution; in case of supersaturation the 
 solid will deposit from the solution, and if the solution is now unsaturated, 
 as the result of the rise in temperature, the crystal will dissolve. 
 
 TABLE OF SOLUBILITY OF 
 (Grams dissolved by 100 cm. 3 
 
 SOLIDS. 
 of water.) 
 
 At o 
 
 Sodium Chloride 35.6 
 
 Sodium Nitrate 72.9 
 
 Sodium Sulphate (hy- 
 
 drated) 5.02 
 
 Potassium Chloride 30.0 
 
 Potassium Nitrate 13.3 
 
 Potassium Sulphate 8.46 
 
 Potassium Bichromate . . 4.9 
 
 Ammonium Chloride. . . . 28.4 
 
 Copper Sulphate 18.2 
 
 Calcium Hydroxide 0.174 
 
 Calcium Sulphate 0.205 
 
 Calcium Chloride 49-59 
 
 C. 
 
 grm. 
 
 
 
 35- 6 3 
 87-5 
 
 grm. 
 
 100 C. 
 39.9 grm. 
 180.0 
 
 (Parts dissolved 
 by 100 parts of 
 80% alcohol 
 at 15.) 
 
 1.22 
 
 2.8 
 
 42.4 
 
 (32.38 C.) (anhydrous) 
 
 i. 3 (sp.gr. .94) 
 
 34.7 
 
 56.6 
 
 29.0 
 
 247.0 
 
 (18 
 
 -) 
 
 10.9 
 
 26.2 
 
 1 3. i 
 
 102.0 
 
 37.28 
 
 72.80 
 
 42.31 
 
 203.22 
 
 0.13 
 
 0.08 
 
 0.208 
 
 .218 
 
 74.00 
 
 ' . I49-9 8 
 
 0.45 (70%) 
 0.4 
 
 0.21 (40%) 
 
 12.0 (abs.8) 
 
 60.00 (abs. 80) 
 
92 EXPERIMENTAL CHEMISTRY. 
 
 Heat of Solution. We have observed that the process of solution is 
 usually accompanied by thermal phenomena, i.e., heat is absorbed or 
 evolved. This is called the heat of solution. It is obvious that for differ- 
 ent amounts of water, this will vary; in order to establish uniformity, 
 the heat of solution is now usually understood to be the amount of heat 
 liberated or absorbed by the solution of i gram-molecule of the sub- 
 stance under consideration, in such a large quantity of water (solvent) 
 that the further addition of water will not yield an additional heat effect. 
 The addition of the first quantity of water to the solute produces a rela- 
 tively larger thermal change than the succeeding additions of equal 
 amounts of the solvent, therefore the use of an unlimited volume of the 
 solvent as suggested above. 
 
 Ostwald defines the "heat of solution" as "the heat which is taken 
 up or given out when a further quantity of salt is dissolved in a solution 
 saturated at a definite temperature." "This quantity of heat," he says, 
 "must not be confused with that which accompanies the solution of a 
 salt in the pure solvent, and which is usually what is measured. In 
 the case of difficulty soluble substances, it is true, the two are not greatly 
 different; but where the substances are soluble in large amounts, they 
 can have not only a different value but even a different sign." Richards, 
 T. W., defines the "heat of solution," of a substance, as that which is 
 developed or absorbed when the final quantity (i mole) of salt which 
 finishes the saturation of an unlimited quantity of water, is added. An 
 erroneous application of this principle has led to many apparent and 
 perplexing contradictions. 
 
 The following tables give the heat of solution of a number of substances 
 in a large quantity of water at i8-2o C. This is designated by the 
 addition of the abbreviation Aq (aqua) to the formula or symbol of the 
 substance: 
 
 HEAT OF SOLUTION OF SOLIDS. 
 (From Hortsmann, Theoret. Chem., p. 502.) 
 NaOH, Aq = +9780 cal. KOH, Aq = +12500 cal. 
 
 NaCl, Aq = 1180 cal. KOH.2H 2 O, Aq - 30 cal. 
 
 ) 3 , Aq = +5640 KC1, Aq = 4440 
 
 EO H 2 O, Aq = 16160 K 2 SO 4 , Aq = 6380 
 
 ) 4 ", Aq = +460 NH 4 C1, Aq = 3880 
 
 o H 2 O, Aq = 18760 LiCl, Aq = +8440 
 
 NaC 2 H 3 O 2 , Aq - +4200 CaCl 2 , Aq = +3258 
 
 Na Br, Aq = 190 AgCl, Aq = 15800 
 
 Na Br^HjO, Aq = 4710 Ag Br, Aq = 20200 
 
 C 12 H 22 O n , Aq = 800 Ag I, Aq = 26600. 
 
 HEAT OF SOLUTION OF LIQUIDS. 
 
 H 2 SO 4 , Aq = +17800 cal. CH 3 OH, Aq = +2000 cal. 
 
 C 2 H 4 O 2 , Aq = +420 cal. C 2 H 5 OH, Aq =: +2540 cal. 
 
 (C 2 H 5 ) 2 O, Aq = +5940 cal. C 3 H 7 OH, Aq = +3050 cal. 
 
SOLUTIONS. 93 
 
 HEAT OF SOLUTION OF GASES. 
 
 H 9 F 9 , Aq = +11800 cal.* Cl.,, Aq = +4870 cal. 
 
 HC1, Aq =: +17310 cal. CO 2 , Aq = +5880 cal. 
 
 HBr, Aq = +19940 cal. NH 3 , Aq = +8430 cal. 
 H I, Aq = +19210 cal. 
 
 All gases which have been examined, and all liquids, as a rule, dis- 
 solve with evolution of heat, but solid substances show no such uniformity 
 of conduct. Some of them, solids, dissolve with an absorption of heat 
 while others behave in a contrary manner. Nernst suggests that the 
 explanation of this is simple. Assume that a given substance in the 
 gaseous condition always dissolves with an evolution of heat, then if it is 
 reduced to the liquid state before it is placed in solution, it will dissolve 
 with an evolution or absorption of heat according as its heat of vapori- 
 zation is less or greater than its heat of solution. The sign of the heat 
 of solution of the substance in the solid phase will depend upon the 
 relative magnitude of its heat of sublimation and its heat of solution 
 when in the gaseous phase. Interpreting the data presented by the above 
 tables in the light of the above suggested explanation, we conclude that 
 the " heat of solution in the gaseous state is always greater than the heat 
 of vaporization, but it is usually smaller than the heat of vaporization 
 plus the heat of fusion, i.e., it is smaller than the heat of sublimation." 
 
 Effect of Pressure on Solubility oj Solids. A solution usually occupies 
 less space than the sum of its components; it never occupies more than 
 the sum. In some cases, for example, solutions of Li(OH), Ba(OH) 2 , 
 NiSO 4 , CaSO 4 , the volume of the solution is less than the original volume 
 of the solvent. Inasmuch as the process of solution is frequently accom- 
 panied by small changes in volume, it is evident that if a solution is in 
 equilibrium with the solid substance, pressure will alter the solubility 
 in a manner which may be anticipated by the application of the principle 
 enunciated in Le Chatelier's Theorem. The effect of pressure is so 
 small that pressures equivalent to hundreds of atmospheres are required 
 to effect a change in solubility so small that it can scarcely be measured. 
 
 The question may arise, what is the explanation of these volume- 
 changes observed while preparing solutions; why does alcohol (ethyl) 
 and water when mixed in equal proportions by volume, contract about 
 3 per cent, of the total volume? Why do 10,000 cm. 3 of water and 27.5 
 cm. 3 (58.5 grm.) of sodium chloride yield a solution whose volume is 
 about 10016.5 cm. 3 ? The question can not as yet be answered definitely. 
 It has been suggested that the "associated" molecules of the water are 
 broken up during the dissolving of the solute, and as the molecules 
 occupy less space when in a " non- associated " condition, this may ac- 
 count for the observed contraction. Again, it may be due to the fact 
 that the dissolved substance undergoes hydration, i.e., attracts water 
 to itself, forming a hydrate. Now the water in hydrated salts 
 occupies less space than when not thus combined. This may partially 
 * 580 cal. (ext. work) should be subtracted from above figures. 
 
94 EXPERIMENTAL CHEMISTRY. 
 
 explain the conduct of the solutions referred to above as examples in 
 which the volume of the solution was less than original volume of solvent. 
 Just where the contraction is no one seems to know. 
 
 Solubility of Liquids. It is frequently convenient, when speaking of 
 the relative solubility of liquids in liquids, to refer to the solubility of 
 " pairs of liquids. " These groups of two many be divided into four orders : 
 
 SOLUBILITY OF PAIRS OF LIQUIDS. 
 
 i. Miscible in all proportions. 
 
 2. Partially miscible. 
 
 3. Immiscible (mutually insoluble). 
 
 4. Pairs which are in one order at some temperatures but in another 
 order at other temperatures. 
 
 The first order embraces those pairs in which the solubility of the 
 two liquids in each other is unlimited; the second order, those pairs in 
 which the solubility of the two liquids in each other is limited; the third 
 order, those pairs in which the two liquids are not wholly but practically 
 insoluble in each other, as mercury and w r ater, organic liquids and water, 
 etc.; the fourth order, explains itself, and suggests that orders, one, two 
 and three, will pass gradually from one to the other with sufficient change 
 of temperature. 
 
 Experiment IX. Relative Solubility of Liquids. 
 
 (a) To 5 cm. 3 of water in a test tube add a few drops of alcohol; shake. 
 How many phases are present? Do alcohol and water mix? Add 
 alcohol in small quantities until 5 cm. 3 of alcohol have been added? 
 Does the mixture become homogeneous on shaking after each addition 
 of alcohol? To which order does the above pair of liquids belong, 
 basing your answer on observed results ? 
 
 (b) Repeat (a) using 5 cm. 3 of ether. 
 
 (c) Repeat (a) using 2 cm. 3 of kerosene. 
 
 (d) Repeat (a), using 2 cm. 3 of benzol or toluol. Define "phase." 
 
 Experiment X. Mutual Solubility of Liquids. 
 
 To 3 cm. 3 of distilled H 2 O in a test tube add 3 cm. 3 of ether and shake 
 vigorously. How many phases are present ? Does each layer represent 
 a solution? Set tube aside to be used later. To 3 cm. 3 of ether add 
 3 cm. 3 of distilled H 2 O. Proceed as above. Compare the contents of 
 the two test tubes. Pour the contents of one tube into the other tube, 
 thereby mixing the solutions; allow tube to stand until there are two 
 distinct phases, i.e., two layers. Dehydrate about a grm. of pulverized 
 CuSO 4 .io H 2 O, the salt should be nearly white. Now devise a method 
 for proving that the top layer of the solution in the test tube is a solution 
 of water in ether. Also prove that ether is present in the lower layer 
 which is a solution of ether in water. 
 
 Hint. When a solution of ether and water is gently heated in a test 
 tube, the ether is rapidly expelled and can be inflamed. 
 
SOLUTIONS. 95 
 
 The heavier liquid at the bottom of the test tube is an aqueous solution 
 of ether containing about 10 per cent, of ether; the upper layer is an 
 etheral solution of water containing about 3 per cent, of water. A 
 similar phenomenon is exhibited by the use of a concentrated solution 
 of potassium carbonate and aqueous ammonia, sp. gr. 0.88. This is 
 probably the only known case in which aqueous solutions of inorganic 
 substances behave in this manner. 
 
 Experiment XI. Effect of Temperature. Critical Solution Tempera- 
 ture. 
 
 (a) Increase in temperature produces complete miscibility. To a test 
 tube containing about 3 cm. 3 of solid or liquid phenol* add about 10 
 cm. s of water; cork the tube and shake vigorously; a milky appearing 
 mixture, an emulsion, is the result; allow the tube to stand for some 
 time, or better, remove the cork and heat gently when the mixture will 
 separate into layers an aqueous solution of phenol above, and a solu- 
 tion of water in phenol below; now place the test tube in the flame 
 and heat gradually until the line of demarcation between the two phases 
 slowly disappears; warm a thermometer until its reading is about 80 C., 
 then introduce it into the test tube so that its bulb will be immersed 
 in the contents, and determine the temperature at which the mixture 
 becomes homogeneous; cool the tube and note the reading of the ther- 
 mometer when the milky appearance is observed. Take the average 
 of the two thermometer readings. What are your conclusions as to the 
 mutual solubility of this pair of liquids at the temperature at which the 
 contents became homogeneous? At the temperature of the laboratory 
 (20 C.)? 
 
 (b) Lowering of the temperature produces complete miscibility. Use 
 dimethylamine and water. (Instructions.) 
 
 The lowest temperature at which pairs of liquids like phenol and 
 water become miscible in all proportions is called the "critical solution 
 temperature." 
 
 Experiment XII. Distribution of Solute Between Two Immiscible 
 Solvents. Law of Distribution. 
 
 To a small flake of sublimed iodine in a test tube add 10 cm. 3 of 
 distilled H 2 O; shake it vigorously for several minutes as the iodine is 
 nearly insoluble and dissolves slowly; pour off the clear solution into a 
 clean test tube. Repeat above operation using an aqueous solution of 
 KI as the solvent. Divide each solution into two portions; to one portion 
 of each solution add 3 cm. 3 of ether; and to the other portion of each, 
 add 3 cm. 3 of carbon disulphide; shake the tubes. Results? In each 
 case, after adding ether or carbon disulphide, how many phases were 
 present in the respective test tubes? Drawing your conclusions from 
 
 * Caution: A very small quantity of phenol (carbolic acid, C6H6O) will attack the 
 flesh. Handle vessel containing it with the utmost care. 
 
96 EXPERIMENTAL CHEMISTRY. 
 
 observed results, arrange the solvents in the order in which the increasing 
 solubility of iodine is shown. 
 
 Law of Distribution. The solute is distributed between the two 
 solvents in such a way that the ratio oj its concentration in each is a con- 
 stant. This constant is practically independent of the absolute con- 
 centration and is dependent only upon temperature and the nature of 
 the solute and the two solvents. (It is obvious that the ratio of the 
 concentrations is the ratio of the solubilities in the separate solvents.) 
 Also, if two or more substances are placed in solution the coefficients of 
 distribution are the same as if each substance were present alone. (This 
 reminds one of the analogous conduct of gases.) In connection with 
 the above experiment, it may be noted that in i cm. 3 of the carbon 
 disulphide solution there will be found 600 times as much iodine as in 
 i cm. 3 of the aqueous solution of iodine. 
 
 Experiment XIII. Pressure of Gas Affects Solubility. 
 
 Half fill the large test tube which you used in generating oxygen, with 
 water; cork the tube with a rubber stopper perforated with a single hole; 
 shake the tube until the water is saturated with air; connect tube with 
 an air pump and exhaust the air from above the solution; observe the 
 small bubbles of gas passing out of the liquid. Explain. What would 
 be the probable effect of increasing the pressure of the gas ? 
 
 Henry's Law. At a constant temperature, a given quantity of the liquid 
 solvent will dissolve weights of a gas which are proportional to the 
 pressure of the gas. This law holds when the gases are only moder- 
 ately soluble. When a mixture of gases dissolves in a liquid, each dis- 
 solves as if it were present alone i.e., each dissolves according to its 
 own partial pressure. In other words, the gas distributes itself between 
 the liquid and gaseous phases according to a constant (K) which is a 
 characteristic of each gas. (Recall the Law of Distribution.) 
 
 Experiment XIV. Effect of Solvent on Solubility. 
 
 Half fill a test tube with water; observe that it is perfectly clear; add 
 5 cm. 3 of alcohol. Explain the liberation of the small bubbles of air. 
 
 SOLUBILITY OF GASES. 
 (o C. 760. mm.) 
 
 Name of gas. i Vol. of H 2 O i Vol. of C 2 H O 
 
 Dissolves, Dissolves, 
 
 Hydrogen 0.02 vols. 0.07 vols. 
 
 Oxygen 0.04 vols. 6.28 vols. 
 
 Nitrogen 0.02 vols. 0.13 vols. 
 
 Carbon dioxide 1.79 vols. 4.32 vols. 
 
 Hydrogen sulphide 4.37 vols. 17.9 vols. 
 
 Ethylene 0.25 vols. 3.6 vols. 
 
 Ammonia 1148.80 vols. 
 
 Hydrochloric acid 505,00 vols. 
 
SOLUTIONS. 97 
 
 Experiment XV. Effect of Temperature on Solubility. 
 
 Half fill a 1000 cm. 3 flask with water; cork it and shake it until water 
 is saturated with air; determine its temperature. Calibrate a 300 cm. 3 
 flask; provide it with a rubber cork through which passes a delivery tube 
 which will dip beneath the surface of the water in the pneumatic trough; 
 fill the 300 cm. 3 flask and delivery tube with the water saturated with 
 air; heat the flask to boiling; collect the air which is liberated in an in- 
 verted eudiometer by water-displacement; read the barometer, determine 
 the temperature of the water in the trough, equalize the water levels 
 and take the reading of the eudiometer. Calculate the volume of gas 
 dissolved by a cm. 3 of water under the conditions which prevailed while 
 preparing the aqueous solution of air. 
 
 Note. A test tube may be used in place of the eudiometer by marking 
 the volume of gas with gummed paper, and then calibrating the tube. 
 
 SOLUBILITY OF GASES. 
 (Vols. of gas which i vol. of H 2 O will dissolve.) 
 
 L/J.O. \Ji gClO V 
 
 Hydrogen. 
 
 V JJUt~J.Jl 4. VW1. VJ 
 
 Oxygen. 
 
 'i j-j-ovy win vj.1* 
 
 Nitrogen. 
 
 Carbon dioxide 
 
 .0215 vol. 
 
 .0411 vol. 
 
 .0203 vol. 
 
 .0179 vol. 
 
 .0206 vol. 
 
 .0362 vol. 
 
 .0179 vol. 
 
 .0144 vol. 
 
 .0198 vol. 
 
 .0325 vol. 
 
 .0160 vol. 
 
 .0118 vol. 
 
 .0185 vol. 
 
 .02.83 vol. 
 
 .0140 vol. 
 
 .0090 vol. 
 
 5 
 10 
 
 20 
 
 The quantity of gas which a liquid is capable of dissolving depends 
 upon, (i) nature of solvent; (2) nature of gas; (3) the temperature of 
 the solvent; (4) the partial pressure of the gas. 
 
 Gases with reference to each other, are regarded as being miscible 
 in all proportions. Inasmuch as solids and liquids directly vaporize 
 into gases, the latter are considered as solvents of the former. 
 
 Experiment XVI. Lowering the Vapor Tension or Elevation of the 
 Boiling Point. 
 
 Determine the boiling point of distilled water proceed as in a former 
 experiment. In this case determine the temperature of the liquid as 
 well as that of the vapor. Record your data. Add a sufficient quantity 
 of CaCl 2 to nearly saturate 20-25 cm. 3 of the distilled water; heat the 
 solution to boiling. Observe, (a) the temperature of the solution; (b) 
 the temperature of the vapor. Record data. Explain. Will vegetables 
 cook more rapidly in a boiling salt solution than in boiling distilled water ? 
 Why? 
 
 Caution. Suspend the thermometer when taking temperatures; do 
 not allow it to touch the flask. 
 
 Experiment XVII. Depression of the Freezing Point. 
 
 Half fill a test tube with distilled water; determine its freezing point; 
 add NaCl until the water is nearly saturated; determine the freezing 
 
 7 
 
9 8 
 
 EXPERIMENTAL CHEMISTRY. 
 
 point of the solution. Has the freezing point of the water been de- 
 pressed? How much (express in degrees C.)? (See Fig. 21, Beckmann 
 apparatus.) 
 
 Elevation of the Boiling Point. It has been known for many years that 
 the normal boiling or freezing point of a liquid may be altered by dissolv- 
 
 ing a quantity of a solid substance in it. We 
 have observed that pure water freezes at o C. 
 and boils at 100 C. (760 mm.), whereas the 
 freezing point of solutions of solids is below, 
 and the boiling point above, normal. If the 
 boiling point of a solvent is elevated, it is 
 obvious that its vapor tension is lowered by the 
 presence of the solute. 
 
 It has been shown experimentally that a 
 definite relation holds between the freezing and 
 boiling points of a liquid and the quantity and 
 nature of the dissolved substance. Further, 
 the degree of the depression of the freezing 
 point or the elevation of the boiling point de- 
 pends upon four factors: (i) the nature of the 
 solute; (2) the quantity of the solute; (3) the 
 nature of the solvent; (4) the quantity of the 
 solvent. To make an application i gram of 
 salt dissolved in 100 grams of acetic acid pro- 
 duces a depression of the freezing point which 
 is different from that produced by dissolving 
 the same quantity of salt in 100 grams of water. 
 One mole of any substance dissolved in 100 
 grams of solvent must always produce a cer- 
 tain definite increase in the boiling point of 
 that solvent, because it produces a definite 
 depression of the vapor tension. (Recall rela- 
 tion of osmotic pressure to vapor pressure.) 
 
 The elevation produced by i gram of sub- 
 stance in 100 grams of solvent is spoken of as 
 the specific elevation (d). It has been found 
 that the specific elevation (d) produced by any 
 substance in any given solvent when multiplied 
 by the molecular weight (M) of the solute, gives 
 a product which is practically a constant (K) 
 for that particular solvent. K represents the 
 
 molecular-increase of the boiling point, i.e., " that due to the solution of 
 i mole of substance in 100 grams of solvent, which must be constant 
 for all substances in the same solvent." 
 
 FIG. 21. 
 
 K = dM, 
 
 - 
 
SOLUTIONS. 99 
 
 One mole of any substance dissolved in 100 grams of ether increases 
 the boiling point of the latter 21.10 C. 
 
 Ostwald (Physico-Chemical Measurements) gives the following 
 values for the constant K which is equal to 100 K. : 
 
 Ethyl ether .............. 2110 Ethyl acetate ......... 2610 
 
 Benzene ................ 2670 Acetone ............. 1670 
 
 Chloroform .............. 3660 Water .............. 520 
 
 Carbon disulphide ........ 2370 Ethylene dibromide . . . 6320 
 
 Acetic acid .............. 2530 Aniline ............. 3220 
 
 Ethyl alcohol ............ 1150 Phenol .............. 3040 
 
 The following expression has been developed for the purpose of 
 ascertaining molecular weights: 
 
 lOOg 
 
 g 
 M - 100 K --- , 
 
 AG 
 
 where G represents the number of grams of solvent; g, the grams of 
 solute; A the elevation of the boiling point in degrees, and M, the molec- 
 ular weight of the solute. The value of K may also be found by a proc- 
 ess of reasoning based upon thermo-dynamical principles. 
 
 Beckman while determining the molecular weight of iodine in ether 
 recorded the following data: 
 
 K = 21.10, g = 2.0579, G = 30.14, A= 0.566; 
 M = 100 K 
 
 substituting in above equations, 
 
 M = 100 K . 
 
 AG 
 
 2.0579 
 
 M = IOO X 2I.IO 
 
 0.566 x 30.14 
 
 M = 254, 
 I 2 = 254- 
 
 Morgan suggests the use of the following simple proportion in cal- 
 culating the molecular weight: 
 
 2.0579 
 21. i : 0.566 :: M grm. per 100 grm. : x 100. 
 
 30.14 
 M = 254. 
 
IOO EXPERIMENTAL CHEMISTRY. 
 
 Depression oj the Freezing Point. Using a similar method of cal- 
 culating experimental data, we are able to ascertain the molecular 
 weight of the solute by observing the depression of the freezing point of 
 the solvent. In 1887 Raoult found that " i mole oj any substance dis- 
 solved in 100 grams oj any one solvent causes a constant depression oj 
 the freezing point." It will be observed that the relation of the solute 
 to the freezing point is very similar to its relation to the boiling point. 
 The mathematical form of the expression used to calculate the molecular 
 weight is identical with the above. If d equals the specific depression 
 and K the molecular depression, then, 
 
 K 
 
 K = dM, M = . 
 d 
 
 AG 
 
 lOOg 
 
 g 
 
 M = 100 K -- , 
 
 AG 
 
 where G represents the number of grams of solvent; g, the grams of 
 solute; A the depression of the freezing point in degrees, and M the 
 molecular weight of the solids. 
 
 SOLVENTS. 
 
 Acetic Acid. Water. 
 
 Solute. d. M. K. Solute. d. M. K. 
 
 SO 2 0.6015 64 38.5 NH 3 1-1705 17- 18.80 
 
 CS 2 0.5050 76 38.4 CuSo 4 0.1153 159.7 18.41 
 
 CH C1 3 0.3247 119.3 38.8 C 2 H 4 O 2 o.3i6 2 60, 18.97 
 
 Aver. 38.6 Aver. 18.72 
 
 (Ostwald), 38.8 (Ostwald), 18.90 
 
 (Jones), 18.50 
 
 SOLUTIONS IN WATER. 
 
 (Jones, Phys. Chem.) 
 
 Solute. Molecular. Solute. Molecular. 
 
 (Organic.) Lowering (K). (Inorganic.) Lowering (K). 
 
 Methy alcohol 17.3 Hydrochloric acid 39.1 
 
 Malic acid 18.7 Sodium oxalate 13.2 
 
 Acetone 17.1 Sodium hydroxide 36.2 
 
 Acetic acid 19.0 Sulphuric acid 38.2 
 
SOLUTIONS. 101 
 
 RAOULT'S LAW FOR SOLVENTS. 
 (Raoult's Tables). 
 
 M K K/M 
 
 Solvents. Mol. Weight. Mol. Lowering. Lowering produced by 
 
 i molecule in 100 mole- 
 
 cules. 
 
 Formic acid 46 29 o-63 
 
 Acetic acid 60 39 -65 
 
 Benzene 78 ,5 o.64 
 
 Ethylene bromide 123 * v '-T5 ' ? '^ -59 
 
 Water 18 , 47 2.6i 
 
 Raoult's Law. "// one mohci'le ' of * "any ~ytel)s&&(?e $s dissolved in 
 one hundred molecules of any liquid of a different nature, the lowering of 
 the freezing point of this liquid is always nearly the same, and approx- 
 imately o.63 " Sci. Memoirs. The above table makes apparent the 
 real meaning of Raoult's Law. It also suggests, that inasmuch as the 
 molecular-lowering of water, which in a number of cases is as great as 
 47 and is an exception to the law, that its molecular weight is probably 
 greater than 18. Ramsay and Shields, from a study of the surface-tension 
 of water, reached the conclusion that .the molecular formula for water in 
 the liquid state should be (H 2 O) 4 or H 8 O 4 . This conclusion is confirmed 
 by the freezing point lowerings of water produced by dissolved substances. 
 
 K 
 
 The value 2 .61, ( ), 
 M 
 
 indicates that the weight of solvent used, contained only one-fourth of the 
 number of molecules required, or in other words, each molecule weighed 
 four times as much as calculated. Assuming that the weight of the mole- 
 cule should be 18 x 4 = 72, and inserting this value in the equation, 
 
 K 47 
 
 = (o.63, Aver.), we have = o.65. 
 
 M 18x4 
 
 This value is very close to the other values of the table. 
 
 It should be remembered that Raoult's Law is only an approximation, 
 but one of much value in correlating many facts. 
 
 Experiment XVIII. Lowering of the Solution Tension. 
 
 Half fill a 250 cm.3 or 500 cm.3 flask with distilled H 2 O; (The flask 
 should have a long narrow neck) saturate the water with ether; add H 2 O 
 and ether repeatedly until flask contains a saturated aqueous solution of 
 ether which rises into the neck of the flask for a distance of about 2 cm. 3 ; 
 there should be a layer of ether about i cm. in thickness above the aqueous 
 solution. The height of each should be marked with pieces of gummed 
 
102 EXPERIMENTAL CHEMISTRY. 
 
 labels. Now add 3 or 4 cm. 3 of benzol, or toluol, either of which is quite 
 soluble in ether; mark its height on neck of flask; shake vigorously for a 
 minute (avoid spilling the contents of the flask) ; allow flask to stand un- 
 disturbed for a few minutes. Observe the height of the aqueous solution. 
 Note the thickness of the layer of ether. Is there more or less ether in 
 solution in the water than there was before adding the benzol to the ether ? 
 Why did the layer of ether increase in thickness by the process ? 
 
 The answer to the question as to whether the process of solution is to 
 be regarded as purely chemical or physical must be delayed until such time 
 as the evidence sh^.ll preponderate for one or the other of the two views, 
 which is certainly not the case with oar present information. The question 
 has been much discussed,, am} jbecause of the alteration in value of well 
 defined properties of a sutstance when it is dissolved in a liquid, there are 
 many chemists who regard solution as a chemical process. As a rule, 
 the sum of the values of a property common to solute and solvent is not 
 the value of this property for the solution. Thermal and volume rela- 
 tions, solutions of constant boiling point, hydration and other phenomena 
 which appear during the dissolving, point to the existence of a certain 
 affinity between solute and solvent. There is little doubt but that the 
 two react and influence each other's properties, but as to the origin and 
 nature of this influence chemists are unable to offer a probable expla- 
 nation. 
 
 Those chemists who maintain that solution should not be regarded as a 
 chemical process point to the conduct of the solute in dilute solutions, with 
 respect to its volume, temperature and pressure relations, i.e., the appli- 
 cability of the gas law. In such cases the influence of the solvent may be 
 neglected altogether. Again, we find that the quantity of a solid which 
 a liquid will dissolve varies with temperature; such variation in the com- 
 position of chemical compounds is unknown. The impossibility of ex- 
 pressing the concentration of saturated solutions in terms of integral multi- 
 ples of the chemical combining weights, and the easy recovery of the solid 
 solute by evaporation, together with the above mentioned phenomena tend 
 to suggest that the process of dissolving is physical. 
 
 The subject of solutions will be discussed in a future chapter. 
 
 PROBLEMS. 
 
 i. How many grm. of KOH will be required to prepare 500 cm. 3 of a 
 i N solution? An .01 N solution? A 5 N solution? 
 
 2. Calculate the relative number of grm. of each solute in normal solu- 
 tions of H 2 SO 4 , HC1, H 3 PO 4 . In 3 N solutions. In 2 N solutions. 
 
 3. How many grm. of HNO 3 in 1000 cm. of a 5 per cent, solution 
 (aqueous) ? If 10 grm. of NaOH are dissolved in sufficient H 2 O to give 
 a 5 per cent, solution of the solute, what was the total volume of the solu- 
 tion? 
 
 4. How many cm. 3 of water will have to be added to 10 grm. of H 2 SO 4 
 (sp. gr. 1.84) to yield a 10 per cent, solution of acid? 
 
SOLUTIONS. 103 
 
 5. If i 1. of H 2 O absorbs i 1. of CO 2 at o, 760 mm. how many grams 
 of CO 2 gas are contained in a bottle of carbonic water holding 200 cm.s 
 of solution, the pressure being 5 atmospheres ? 
 
 6. The approximate composition of air is 30.9 percent, of O, and 79.1 
 of N, by volume. At 15 water absorbs 0.0299 volumes of O and 0.0148 
 of N, the pressure of each being that of the atmosphere. What is the 
 composition of air absorbed in H 2 O ? 
 
 Ans. By volume, 34.8 per cent, of O and 65.2 per cent, of N. 
 
 7. The increase in the boiling-point of 54.65 grm. of CS 2 caused by 
 the addition of 1.4475 gr m - of P is o.486. What is the molecular weight 
 of P in CS 2 ? What is the molecular formula if the atomic weight is 31 ? 
 
 8. The molecular weight of a substance is 60. If 10 grm. of this 
 substance is dissolved in 100 grm. of a solvent, the increase in the boiling 
 point is o.87. Calculate the molecular increase in the boiling point. 
 
 9. Find the molecular weight of oxime, (CH 3 ) 2 CNOH, if 0.284 grm. 
 of it causes a decrease of o.i55 in the freezing point of 100 grams of 
 glacial acetic acid. 
 
 10. A student working in this laboratory tabulated the following 
 data while determining the molecular weight of turpentine by the " boiling 
 point method." Boiling point of pure solvent (ether) as indicated by a 
 Beckmann thermometer, 3. 25; weight of solvent 18.45 g rm -5 weight of 
 turpentine, 0.865 g rm -; boiling point of solution, 3-96. Calculate the 
 molecular weight of ether. 
 
 n. The following data was recorded while determining the molecular 
 weight of sugar by "freezing point method." 
 
 Freezing point of pure H 2 O, (Beckmann) 5.475, 
 
 Freezing point of solution after adding i gram of sugar to 20.09 
 grams of H 2 O, 5. 198. Calculate the molecular weight of sugar. 
 
CHAPTER XII. 
 ACIDS, BASES AND SALTS. 
 
 NOMENCLATURE. 
 CHLORINE. 
 
 Symbol, Cl. At. Wt. 35.45. 
 
 Chlorine is an elementary substance belonging to the class of " non- 
 metals." Under ordinary laboratory conditions it is a heavy, greenish- 
 yellow gas which possesses an irritating odor and poisonous properties. 
 It was discovered by Scheele in 1774, who prepared it by heating hydro- 
 chloric acid with the oxide of manganese. Davy, however, proved its ele- 
 mentary nature and gave it its name in 1810. It possesses a great deal of 
 chemical energy, i.e., shows great tendency to react chemically with other 
 substances, therefore, it is not found in nature as jree chlorine. Observe 
 during the performance of the following experiment whether chlorine 
 possesses the physical properties characteristic of metals. 
 
 Experiment I. Preparation of Chlorine. 
 
 Arrange a 300 cm. 3 flask (Fig. 22) provided with funnel tube and 
 a delivery tube, so that it connects with washing bottle which should 
 be about one-fourth full of concentrated sulphuric acid. The flask 
 should be placed on a piece of iron gauze resting on a ring clamp or tripod 
 so that flask may be heated. Introduce into flask 20 grams of manganese 
 dioxide (MnO 2 ), preferably in the powdered condition. Place six clean 
 dry bottles of about 150 cm. 3 capacity on your desk and cover each with 
 a glass plate or a sheet of heavy damp paper; procure from the assistant 
 a little powdered antimony on a watch glass, a strip of bright colored calico, 
 some litmus paper, write your name with ink upon a piece of paper, 4 or 5 
 cm. 3 of turpentine in a test tube and a test tube half filled with distilled 
 water; then arrange your hydrogen generator so that you can provide 
 a hydrogen flame on short notice. Now add by means of funnel tube, 
 sufficient concentrated hydrochloric acid to cover the MnO 2 in flask; be 
 sure that lower end of funnel tube is dipping into acid; heat flask very 
 gently; by means of a delivery tube bent at right angles and attached 
 to the wash bottle so that it can be inserted into the mouth of a bottle, 
 collect gas by vertical displacement of air. When bottle presents a 
 greenish-yellow appearance remove it and cover its mouth with glass plate 
 or damp paper. Repeat operation until six bottles have been filled with 
 the gas and water in test tube has been saturated with chlorine. Then 
 
 IQ\ 
 
ACIDS, BASES AND SALTS. 105 
 
 proceed with the following experiments. The equation for the above 
 reaction may be written as follows: 
 
 MnO 2 + 4 HC1 MnCl 4 + C1 2 + 2H 2 O. 
 MnCl 4 MnCl 2 + C1 2 . 
 
 Is chlorine lighter or heavier than air? If a liter of hydrogen weighs 
 .0899 grams under standard conditions what will the same volume of 
 chlorine gas w r eigh under identical conditions ? 
 
 Experiment II. Properties of Chlorine. 
 
 (a) Bleaching properties. Into a bottle of gas insert pieces of dry 
 calico, litmus paper and paper with carmine ink upon it; set it aside for 
 fifteen or twenty minutes. Repeat above operation with another bottle 
 of chlorine, but moisten the articles before inserting them. Examine be- 
 
 FIG. 22. 
 
 fore leaving laboratory. Results? Compare with result when articles 
 were perfectly dry. What conclusion would you draw? It should be 
 observed that organic material is the basis of the coloring matter present 
 in above specimens. 
 
 (b) Combination of antimony and chlorine. Into a bottle of the gas, 
 drop a pinch of powdered antimony. Result? Place another portion 
 of the powder in a test tube and heat it until quite warm, then quickly 
 introduce it into the same bottle of gas. Result ? What factor induced 
 the reaction ? The valency of antimony is three, that of chlorine is one. 
 Write the equation for reaction. 
 
 (c) Affinity of chlorine and hydrogen. 
 
 Insert "hydrogen flame" into a bottle of chlorine (Fig. 23). Result? 
 
106 EXPERIMENTAL CHEMISTRY. 
 
 Hold a piece of moist blue litmus paper just outside of mouth of bottle 
 while " flame" is present on inside. Result ? Hold another piece of blue 
 litmus paper near the mouth of the reagent bottle labeled " hydrochloric 
 acid." Results? What are your conclusions as to the gas formed by 
 the burning of chlorine in hydrogen ? Write equation for reaction. 
 
 Was light and heat evolved as the result of the reaction ? What was 
 probably the main source of this energy? Was it an endo- or exother- 
 mic reaction? Interpret following equation: 
 
 H, (i gram) + Cl, (35.45 grams) > HC1, (36.45 grams) + 22,000 cals. 
 (d) Chemical action on turpentine (C 10 H 16 ). 
 
 Place a strip of filter paper in the test tube containing the C 10 H 16 and 
 heat tube gently until it is warm. Turpentine is very inflammable. In- 
 troduce paper saturated with warm turpen- 
 tine into a bottle of chlorine gas. Result ? 
 If you were told that the black appear- 
 ance of paper was due to free carbon (soot) 
 what would be your conclusion ? Equation ? 
 (e) Substituting power of chlorine, 
 i. Make a few strips of "test paper" by 
 dipping pieces of filter paper in a solution 
 of potassium iodide (KI) with which a 
 clear solution of starch has been mixed. 
 Free iodine will impart a blue or bluish- 
 FIG. 23. black color to starch. Upon a strip of test 
 
 paper pour a few drops of the water saturated 
 with chlorine. Results? Conclusions? Equation? 
 
 2. Pass hydrogen sulphide (H 2 S) through a water solution of chlorine 
 and observe effect. Before passing the H 2 S into the solution notice the 
 action of the latter upon litmus paper, both red and blue. Compare the 
 result with a similar test made after the H 2 S has been passed into it. In 
 view of observed data, how would you explain results ? Equation ? 
 
 3. Optional. (L. T.) Fill a test tube with a saturated solution of 
 chlorine; invert it mouth under water; place it where direct sunlight 
 may fall upon the tube; examine after a day or two. Results? Test 
 water of tube with litmus paper. Results? If any gas has accumu- 
 lated in upper portion of tube, test it with a glowing splinter. Results ? 
 What are your conclusions as to reaction ? Equation ? 
 
 4. Optional. Introduce into a bottle of chlorine a piece of metallic 
 sodium heated gently in a dephlegrating spoon. Result? Equation? 
 
 Experiment III. Optional Methods for Preparation of Chlorine. 
 (Oxidation Processes.) 
 
 (a) Use equal parts by weight of NaCl and MnO 2 ; add strong sulphu- 
 ric acid. 
 
 Mn0 2 + NaCl + tt^O, MnSO 4 + Na,SO 4 + 2H 2 O + 2 CL 
 
ACIDS, BASES AND SALTS. 107 
 
 For stages of reaction (Instruction). 
 
 (b) Potassium permanganate, with hydrochloric acid added slowly 
 from a "dropping" funnel. 
 
 KMnO 4 + 8HC1 - KC1 + MnCl 2 + 4H 2 O + sCl. 
 
 (c) To a few crystals of potassium chlorate, add concentrated hydro- 
 chloric acid. (This is a convenient and efficient method for preparing 
 small quantities of the gas. It is frequently resorted to by analytical 
 chemists.) The reaction may be represented by the following equation: 
 
 2 KC1O 3 + 4HC1 2KC1 + C1 2 O 4 + 2H 3 O + 2C1. 
 
 It must be evident to the student that chlorine possesses considerable 
 chemical energy, and that it manifests a remarkable affinity for hydrogen 
 the majority of the foregoing phenomena being explained by this fact. 
 Its bleaching properties in the presence of moisture appear to be due to 
 the reaction whereby the hydrogen is withdrawn from the water, liberat- 
 ing nascent oxygen which oxidizes the coloring compounds, thus form- 
 ing new substances which are colorless. Chlorine has frequently been 
 called an "oxidizer" on account of this type of reaction. Frequently 
 compounds are " broken up," i.e., suffer decomposition because of the 
 withdrawal of hydrogen by chlorine. 
 
 That which is of dominant interest to us at this time is the compound 
 which is formed as the result of the combination of hydrogen and chlo- 
 rine, and which in the light of former rules of nomenclature, might be 
 termed hydrogen chloride. The properties of this latter substance re- 
 sembled those of hydrochloric acid. This leads us to a study of a class 
 of compounds called "acids." 
 
 ACIDS. 
 
 In the study of the element chlorine which is a typical non-metal, 
 a compound containing chlorine and hydrogen was prepared, which in a 
 number of its properties resembled a class of compounds called acids. 
 The investigation of the properties of this compound may warrant the 
 making of certain generalizations which will be of service to us. 
 
 Experiment I (L.T.). Electrolysis or Analysis of Hydrochloric Acid. 
 
 Fill a Hoffman V-tube for the electrolysis of hydrochloric acid with a 
 dilute solution of the shelf reagent. Connect the positive pole of the 
 battery with the end of the V-tube which is not ordinarily sealed, but is 
 closed with a cork through which a platinum electrode passes. Pass the 
 current; hold a piece of "test paper" near the open end of tube for three 
 or four minutes. Results ? Conclusions ? 
 
 Test the gas collecting in the other arm of the tube by bringing a lighted 
 match near it. Results ? Identity of gas ? What are your conclusions 
 as to the composition of hydrochloric acid ? Is your view supported by 
 any of the experiments performed with chlorine? Explain. At which 
 
io8 
 
 EXPERIMENTAL CHEMISTRY. 
 
 electrode was free chlorine (non-metal) liberated ? This is a characteristic 
 of non-metals. 
 
 Experiment II. Volumetric Composition of Hydrochloric Acid. 
 (a) Fill the Hoffman apparatus which was used for the electrolysis of 
 water, with a solution of hydrochloric acid. Pass the current for fifteen 
 or twenty minutes simply to saturate the water with hydrogen and chlo- 
 rine gases, as the latter is particularly soluble; open stop cocks, forcing 
 out the gases; now proceed with the electrolysis to determine 
 the ratio of the combining volumes of hydrogen and 
 chlorine. Recalling Avogadro's rule, what would you say 
 in regard to the relative number of atoms of each gas in a 
 molecule of the acid? 
 
 (b) Optional. Relative volumes of hydrogen and 
 chlorine in hydrogen chloride. 
 
 Generate hydrogen chloride as suggested in above ex- 
 periments, then dry gas by bubbling it through a w r ash 
 bottle filled with sulphuric acid, or generate it by allowing 
 concentrated sulphuric acid to drop slowly from a separa- 
 tory funnel into a flask containing concentrated hydro- 
 chloric acid. Fill a colorimetric tube or eudiometer (Fig. 
 24) with the dry gas. Introduce about 10 grams of sodium 
 amalgam and quickly cork the open end. Raise and low r er 
 either end of the tube so that the amalgam may come into 
 contact with all portions of the gas. Uncork the tube 
 under mercury. Result ? Adjust, so that liquid on out- 
 side and inside of tube are at same level. What is the 
 volume ? Place thumb over open end of tube and lift it 
 from the liquid; invert; remove thumb and apply a lighted 
 FIG. 24. match. Results? What gas was it? 
 
 What are your conclusions as to the relative volumes of 
 the two constituents of hydrogen chloride ? As to the volume of either 
 constituent relative to the total volume ? 
 
 Experiment III. Preparation and Properties of Hydrogen Chloride Gas. 
 
 (a) Place a small quantity of ammonium chloride in a test tube; add 
 a few drops of concentrated sulphuric acid; hold a piece of blue litmus 
 near mouth of tube. Results ? Repeat with red litmus paper. Results ? 
 Blow your breath across mouth of tube. Results? Pour a few cm. 3 
 of ammonium hydroxide into a test tube, then bring its mouth near to 
 mouth of tube in which gas is being generated. Results? 
 
 2 NH 4 C1 + H 2 SO 4 (NH 4 ) 2 SO 4 -f THC1. 
 
 1 mol. grm. wt. -2 mol. grtn. wts. 
 
 2 mol. grm. wts. i mol. grm. wt. 
 
 To what class of chemical reactions does the above reaction belong? 
 Write equation for the reaction that occurred when the two test tubes 
 with contents were brought close together. 
 
ACIDS, BASES AND SALTS. 1 09 
 
 Experiment IV. Study of the Properties of the Gas Evolved from Hy- 
 drochloric Acid. 
 
 Now half fill a test tube with the shelf reagent known as " hydrochloric 
 acid." Using this, repeat the above experiment in its various parts, and 
 record results. 
 
 Does there seem to be any similarity in the properties of hydrogen 
 chloride gas and the gas evolved from hydrochloric acid ? 
 
 Experiment V. Laboratory Method for Preparation of Hydrogen Chlor- 
 ide Gas. 
 
 Arrange apparatus as you did for the preparation of chlorine with this 
 exception, that you raise the tube of the wash bottle which ordinarily 
 dips into the washing fluid (water in this experiment) so that its lower end 
 just escapes the surface of the water. Pour 15 cm. 3 of water into a beaker; 
 add slowly and while constantly stirring 35 cm. 3 of concentrated sulphuric 
 acid; let the mixture cool. Introduce 35 grams of sodium chloride into 
 generating flask and observe that all parts of system of apparatus are prop- 
 erly arranged. Place upon the desk a f ew r dry bottles, and a label cut into 
 narrow strips about a cm. long. Set a dry bottle under delivery tube of wash 
 bottle to collect gas by vertical displacement of air. Add slowly through 
 funnel tube about one-half of cooled mixture of water and acid, if neces- 
 sary add all of it, and heat gently. Fill three bottles with the hydrogen 
 chloride gas. Half fill a clean test tube with distilled water; mark its 
 height by means of a piece of label; permit the gas to bubble into water in 
 tube until it is saturated; this test tube should be placed in a bath which 
 is at a temperature as near o C. as can conveniently be arranged; at 
 least, place tube in a bottle filled with cold water. While waiting for 
 water to become saturated, proceed with following experiments. The 
 above reaction may be represented as follows: 
 
 NaCl + H 2 SO 4 NaHSO 4 + HCL 
 NaHSO 4 + NaCl - Na^SO, + HCL 
 
 Experiment VI. Solubility of Hydrogen Chloride Gas. 
 
 Invert one of the bottles of gas over a pneumatic trough, with its mouth 
 under water. Results ? Conclusions ? Proofs ? 
 
 Experiment VII. Density of Hydrogen Chloride Gas. 
 
 Using another bottle of the gas, devise a method which will prove 
 approximately that the gas is heavier than air. Can gas be " poured ?" 
 
 Experiment VIII. Properties of Hydrogen Chloride Gas in Solution 
 and Hydrochloric Acid. 
 
 Before proceeding with this experiment, take up Exp. IX. Divide 
 the aqueous solution of the gas into three equal parts. 
 
 (a) To one part in a test tube add a piece of granulated zinc. Results ? 
 Hold a lighted match near the mouth of the test tube. Result ? Con- 
 
110 EXPERIMENTAL CHEMISTRY. 
 
 elusions? Repeat this, and all of the following experiments, by using 
 hydrochloric acid from the shelf. Results ? 
 
 (b) To the second part of the solution add a little sodium carbonate. 
 Results ? 
 
 (c) Divide the third part into four portions and dilute each with twice 
 its volume of water. To one of the portions add a few drops of silver 
 nitrate. Result? To another portion add a few drops of mercurous 
 nitrate. Results ? To the third portion add lead nitrate. Results ? 
 Divide the fourth portion into two parts. To one part add sodium hy- 
 droxide. Results ? 
 
 To the second part add three times its volume of water. Taste it by 
 placing a drop or two on the -tongue by means of stirring rod. Does it 
 possess a caustic, lye-like, or an acid, sour taste ? 
 
 Dip pieces of blue and red litmus paper into the solution. Results? 
 
 What are your conclusions as to the relation between hydrogen chloride 
 gas in aqueous solution and the shelf reagent, hydrochloric acid? 
 
 Experiment IX. (Quant.) Optional. Density of a Solution of Hydrogen 
 Chloride Gas (Hydrochloric Acid). 
 
 When liquid in above test tube (Exp. V) is apparently saturated with gas, 
 remove it and note volume of contents. Results ? Now carefully remove 
 all traces of label and wipe dry outside and inside of tube above solution; 
 weigh tube and contents. Record weight; mark height of solution with 
 label; pour contents of tube into another tube to be reserved for preceding 
 experiment; wash tube and fill with distilled water to top of "label;" 
 remove label and wipe tube as before; weigh empty tube when clean and 
 dry. Weight ? Find density of solution. Record all calculations. 
 
 Experiment X. General Properties of Acids. 
 
 Make very dilute solutions of the following acids, nitric, sulphuric, and 
 acetic. Taste a drop of each. Test them with both blue and red lit- 
 mus paper. Place a piece of granulated zinc in a dilute solution of each. 
 Test each acid with a solution of sodium carbonate. Make a record of 
 all observations. Do all of the acids have some properties in common ? 
 Write the formulae for hydrochloric, nitric, sulphuric, and acetic acids. 
 What element is possessed in common by all acids? Would you sus- 
 pect a causal relation between the element and the properties possessed 
 in common by acids ? 
 
 Experiment XI. Each Acid Has Characteristic Properties. 
 
 Procure from side shelf reagents 4 or 5 cm. 3 of lead nitrate in a test 
 tube, and an equal volume of barium chloride in another tube. Divide 
 the lead nitrate into two parts. To one part add a few drops of dilute 
 hydrochloric acid; to the other add dilute nitric acid. Record results. 
 
 Repeat the above using nitric acid and sulphuric acid and the barium 
 chloride solution. Record results. Would you say that each acid pos- 
 
ACIDS, BASES AND SALTS. Ill 
 
 sesses properties peculiar to itself in addition to its general acid prop- 
 erties ? 
 
 Write equations for all reactions, underscoring the substance which 
 separated out in the solid form (precipitate). Is the precipitate in equilib- 
 rium with molecules in solution? 
 
 It is difficult to accurately define an acid at this stage of our work. 
 From the foregoing experiments it may be deduced that acids are com- 
 pounds which may be either solids, liquids, or gases. They possess 
 the power to alter certain vegetable colors, namely, to turn blue litmus 
 red; to act upon metals, which displace hydrogen; they have a sour taste 
 and usually contain a non-metal (an electro-negative element) united with 
 hydrogen or hydrogen and oxygen. It is generally conceded that it is 
 the replacable hydrogen of these compounds when in aqueous solution 
 which give them their general acid properties. 
 
 Hydrascids. (HC1,H 2 S). 
 ACIDS, 
 
 Oxascids. (HC10 3 ,H 2 S0 4 ). 
 
 Organic. (HC 2 H 3 O 2 , H 2 C 4 H 4 O 6 ). 
 
 ACIDS, 
 
 Inorganic. (HNO 3 , H 3 PO 4 ). 
 
 Oxascids are sometimes looked upon as non-metallic oxides combined 
 with water. The oxide is called the anhydride (H 2 O.SO 3 ,H 2 O.N 2 O 5 ). 
 
 The part of an acid formula which remains after the hydrogen is re- 
 moved is sometimes spoken of as the acid radical. 
 
 Hypo-ous. (HC1O). 
 
 -ous. (HC10 2 ). 
 OXASCIDS, 
 
 -ic. (HC10 S ). 
 
 per-ic. (HC10 4 ). 
 
 The above nomenclature is a type of that which is resorted to in order 
 to distinguish between. a number of compounds closely related in com- 
 position. The endings indicate the relative degrees of oxidation. 
 
 Binary Compounds are those which are composed of two elements, 
 M 2 O. The names of such substances end .in ide. This rule relative to 
 nomenclature has preference over any other rule. Ternary Compounds 
 are those composed of three elements, as, H 2 SO 4 . 
 
 SODIUM. 
 
 Symbol, Na. At. Wt. 23.05. 
 
 Sodium is a typical metal. In dividing the elementary substances into 
 metals and non-metals, it should be remembered, and as we shall see later, 
 the line of demarcation is nowhere distinctly drawn. Midway between 
 
112 EXPERIMENTAL CHEMISTRY. 
 
 these proposed classes there are such elements as arsenic, antimony and 
 bismuth, whose chemical and physical properties permit, under varying 
 conditions, a classification with either the metals or non-metals. These 
 elementary substances occupying intermediate ground are sometimes 
 called metalloids. It is better to regard the elements as constituting a 
 series with a regular gradation of properties. 
 
 Sir Humphry Davy succeeded in preparing small quantities of sodium 
 by electrolysis of fused sodium hydroxide, about 1807. 
 
 Experiment I. Properties of Sodium. 
 
 (Metallic sodium is usually kept under kerosene. It should always be 
 handled by means of dry forceps.) 
 
 (a) Place a piece of sodium upon a dry paper and cut off a piece the 
 size of a small pea. Is the sodium hard? Observe the color, luster, 
 and the effect of air. Half fill the pneumatic trough with water which 
 does not turn red litmus paper to blue. Drop the small piece of metal 
 upon the water and step back from the trough. Does the metal float? 
 Is there evidence of chemical action taking place ? Fasten a match to a 
 rod; light match, and apply flame to piece of metal. Is there evidence 
 of a flame around the sodium? A flame implies the existence of a gas. 
 Where does the gas come from? 
 
 (b) Wrap a piece of sodium in a little paper; place it in a wire gauze 
 basket; hold it under water and collect the gas which escapes in a test tube 
 by displacement of water; apply a lighted match to mouth of tube. 
 Results ? What gas ? The sodium has apparently displaced what from 
 the molecule of water? 
 
 (c) Test the water in the trough with red litmus paper. Result? 
 Take a little of the water between the fingers. 
 
 Pour 3 cm.3 of the shelf reagent labeled sodium hydroxide into a test tube 
 and dilute with an equal volume of water. Repeat above tests. Would 
 you say that they are identical substances ? If so how much of the hydro- 
 gen was displaced from the molecule of w r ater ? Write equation to repre- 
 sent the reaction of sodium and water. What is the valency of sodium ? 
 
 (d) Recall the Exp. in which metallic sodium was introduced into 
 chlorine. Write equation. 
 
 (e) Flame color of sodium compounds. Dip a platinum wire into a 
 little hydrochloric acid; hold it in flame until it imparts no color to the 
 latter; dip wire into a little sodium hydroxide and hold wire in flame. 
 Is the flame colored ? What is the color ? Repeat all of above operations 
 using the water of the pneumatic trough. Results? Conclusions? 
 
 Sodium is a silver-white metal; it loses its metallic lustre on exposure to 
 damp air due to the fact that under such conditions it readily oxidizes; 
 it is quite soft at laboratory temperatures, but hard at 20 C.; melting 
 point, 96.5 C.; specific gravity, .97; conductor of heat and electricity; is 
 an electro-positive element, separating out at negative electrode, but imme- 
 diately reacts on the water forming a compound known as sodium hy- 
 
ACIDS, BASES AND SALTS. 113 
 
 droxide and liberating hydrogen instead; it possesses relatively a great 
 amount of chemical energy, combining so readily with substances that 
 it is not found in the free state in nature. We have seen it displace an 
 hydrogen atom from the molecule of water and combine with the hydroxyl 
 (OH), yielding a compound known generically as a base or an alkali, and 
 specifically, as sodium hydroxide or sodium hydrate, which has the property 
 of turning red litmus paper blue. 
 
 BASES. 
 
 Experiment I. General Properties of Bases. 
 
 Place 3 cm. 3 of each of the following in different test tubes, and dilute 
 each with four times its volume of water: sodium hydroxide, potassium 
 hydroxide, ammonium hydroxide and calcium hydroxide. Test each by 
 means of red and blue litmus paper, tasting a drop, and by taking a little 
 between the fingers. Tabulate data. What are your conclusions as to 
 their similarity with respect to the properties of changing red litmus to 
 blue litmus, alkaline taste (alkalinity, causticity), and lye-like or soapy 
 feeling to touch ? 
 
 Write the formula of each. Is the hydroxyl common to all ? Would 
 you infer a causal relation between properties in common and constituents 
 in common ? Ask the instructor to show solid NaOH and KOH, to you. 
 
 Experiment II. Each Base Has Characteristic Properties. 
 
 To a few cm. 3 of lead nitrate solution add a few cm. 3 of dilute hydro- 
 chloric acid. Divide the precipitate into two parts. To one part add 
 considerable NaOH (reagent). Result? To the other portion, add 
 NH 4 OH in excess. Result? 
 
 Using a solution of silver nitrate repeat above experiments. Results ? 
 Write equations for above reactions. 
 
 As the result of the combination of sodium and the hydroxyl (OH) 
 we have a compound which is typical of a large group of compounds which 
 are known collectively as bases. They seem to possess properties which 
 are the very opposite of those possessed by acids. Bases have the 
 power to change red litmus to blue; they have an alkaline or bitter taste 
 and feel soapy to the touch. The term base has been very loosely, and 
 frequently, inaccurately applied. It is now usually restricted to the 
 hydroxides or hydrates of the metals. It seems to be quite well established 
 that the basic or alkaline properties of the bases in aqueous solution, are 
 due to the action of the hydroxyl. Our concept of a base, in regard to its 
 constituents is, that it is a metallic element (metal) combined with one or 
 more hydroxyls, the number depending on the valency of the metal. As 
 would be expected the above enumerated properties are possessed in 
 varying degrees of intensity by the different bases. 
 
 Strong bases, like sodium hydroxide and potassium hydroxide, are fre- 
 8 
 
114 EXPERIMENTAL CHEMISTRY. 
 
 quently called alkalies. Again, sodium hydroxide and potassium hy- 
 droxide are respectively, spoken of, as caustic soda and caustic potash 
 
 Vegetable. Fixed. 
 
 Alkalies Alkalies 
 
 Mineral. Volatile. 
 
 The bases are commonly distinguished from one another by using the 
 name of the metal before the word hydroxide. 
 
 Bases are not infrequently considered as being composed of a metallic 
 oxide and water. The oxide is called a basic oxide. BaO +H 2 O 
 Ba (OH) 2 . 
 
 The question might logically be raised at this time, what will occur if 
 an acid and a base (their properties are the opposite of one another) are 
 brought together? 
 
 NEUTRALIZATION SALTS. 
 
 Experiment I. Electrolysis of a Salt. 
 
 Make a strong solution of sodium sulphate. Test the solution with 
 both red and blue litmus paper. Is the solution neutral ? Place solution 
 in U-tube for electrolysis after having added a sufficient quantity of lit- 
 mus solution to give a decided blue color to the entire volume of liquid. 
 What changes of color occur when current is passed ? What is liberated 
 at either pole ? Write equations to represent all reactions. 
 
 Experiment II. A Quantitative Study of the Interaction of Acids and 
 Bases. 
 
 Assemble two burettes (Fig. 25) and clean them thoroughly. Clamp 
 them into proper position. Fill the burette at your right with dilute hydro- 
 chloric acid (i of acid to 20 of water). Fill the left burette with a dilute 
 solution of sodium hydroxide (i part of shelf reagent to 10 of water). 
 Run out acid and alkali so that both burettes give zero reading. Clean 
 a small Erlenmeyer flask; place a piece of white paper under either 
 burette. Place flask under " acid" burette and run into it 9 cm. 3 of acid. 
 Be sure to read from the lower side of meniscus. Add a few drops of 
 litmus solution or a small piece of red litmus paper. Now place flask 
 under "alkali" burette and introduce cautiously enough of the alkali to 
 just turn the litmus color to a permanent blue. From time to time shake 
 the flask vigorously to bring alkali and acid into intimate contact, or stir 
 w r ith a glass rod, but don't take the rod out of the flask. Draw from the 
 burettes alternately as is necessary to bring the solution to the point when a 
 drop of either will cause a change in color. When this point has been 
 reached the solution is neutral. Record the number of cm. 3 of each used. 
 The process is known as neutralization. Taste the solution. Place 15 or 
 20 cm. 3 in an evaporating dish and evaporate to dryness. Taste the salt. 
 Has it a familiar taste ? While waiting for the solution to evaporate, re- 
 peat first part of operation by finding how many cm. 3 of the alkali will be 
 
ACIDS, BASES AND SALTS. 115 
 
 required to neutralize 15 cm.s of acid; then repeat again, using 12 cm.3 of 
 acid. Tabulate all data. 
 
 Does the Law of Definite Proportions receive a verification by your 
 results ? 
 
 FIG. 25. 
 
 If time permits, other acids and bases may be used and their ratios 
 determined. Write equations. 
 
 What do you have in the acid that you do not have in the neutral salt ? 
 Answer the same question in regard to the alkali. What has become of 
 these two substances ? Is it possible that their removal from the neutral 
 substances is in any way related to the property of "neutrality?" 
 
1 1 6 EXPERIMENTAL CHEMISTRY. 
 
 Write the equations representing the respective reactions of H 2 SO 4 , 
 HC1, and HN0 3 with KOH, NH 4 OH, and Ca(OH) 2 . 
 
 Experiment III. Interaction of a Metallic Oxide and an Acid. 
 ZnO + 2HC1 ZnCl 2 + H 2 O. 
 
 Experiment IV. Interaction of Metals and Acids. 
 Zn + H 2 S0 4 -*ZnS0 4 + H,. 
 
 Experiment V. Action of an Hydroxide as a Base and as an Acid. 
 
 To a few cm.3 of Pb (NO 3 ) 2 add a few cm.3 of NaOH. Note the pre- 
 cipitate. Filter. Divide the precipitate into two parts. To one part 
 add HNO 3 ; to the other, add an excess of NaOH. Is there any similarity 
 in the action of the acid and the base in these latter reactions ? Write 
 equations. 
 
 Experiment VI. Action of a Base on a Salt. 
 
 To 3 cm.3 of Fe SO 4 add a few cm.3 o f NH 4 OH. Results? Equa- 
 tion? 
 
 Experiment VII. Action of an Acidic Oxide upon a Base. 
 
 CO 2 + Ca (OH) 2 -*CaCO 3 + H 2 O. 
 Experiment VIII. Interaction of Salts. 
 
 HgNO 3 + Na Cl -> HgCl + NaNO 3 . 
 
 Experiment IX. Preparation of a Basic Salt. 
 
 Recall the experiment in which BiCl 3 which is soluble, was changed into 
 an insoluble basic salt by hydrolysis. Define Hydrolysis. 
 
 H 2 -> H- + OH' 
 BiCl s + 2H- + 2OH' Bi(OH) 2 Cl + 2HC1 
 
 1 
 BiOCl + 1^0. 
 
 Experiment X. Preparation of an Acidic Salt. 
 
 Fill a burette with NaOH as found in the reagent bottles. Fill another 
 burette with a cold solution of H 2 SO 4 (i of acid to 2 of H 2 O). Neutralize 
 5 cm. 3 of acid with NaOH. Record number of cm. 3 used. Evaporate 
 solution until it becomes saturated. Crystals will separate out when it 
 becomes cold. To the same volume of NaOH as used above, add twice 
 the quantity of H2SO 4 necessary to neutralize it. Proceed as above. It 
 may be necessary to let solutions stand a couple of days before crystals 
 appear. Compare crystals as to appearance, water of hydration, and 
 reaction toward litmus. Write equations. 
 
 Experiment XI. Classification of Salts. 
 
 Under the heads of Normal, Acidic, and Basic, arrange the following 
 salts with reference to their action toward litmus paper. A few cm. 3 
 
ACIDS, BASES AND SALTS. 117 
 
 of a solution of each will be sufficient. Write formula for each salt. 
 Sodium chloride, sodium hydrogen sulphate, copper sulphate, aluminum 
 chloride, sodium hydrogen carbonate, sodium carbonate and potassium 
 sulphate. 
 
 A salt is the chief product of the interaction of an acid and a base, or (a) 
 it is an acid in which the hydrogen, wholly or in part, has been replaced 
 by a metal, or (b) it is a base in which the hydroxyl has been replaced 
 wholly or in part by the acid radical. Both (a) and (b) are necessary to 
 cover all cases. 
 
 This interaction of acids and bases is a species of reaction known as 
 neutralization. It may be considered as a process in which the removal 
 from solution of the hydrogen and hydroxyls is effected, wholly or in part, 
 by arranging the conditions favorable to their combining and forming 
 neutral water. A more specific detailed definition may be offered later. 
 Neutralization is a quantitative process. 
 
 Would you expect a definite thermal conduct when given quantities 
 of hydrogen and hydroxyls combine to form molecules of water? 
 
 Thermo-chemistry of neutralization. 
 
 NaOH + HC1 NaCl + H 2 O + 137000 cal. 
 KOH + HNO 3 KNO 3 + H 2 O + 137000 cal. 
 Ca(OH) 2 + H 2 SO 4 CaSO 4 + H 2 O + (2 x 137000) cal. 
 
 The heat liberated during the process of neutralization is called the 
 "heat of neutralization." 
 
 Acid Salts. 
 Partial, 
 
 Basic Salts. 
 Neutralization, 
 
 Complete, Normal Salts. 
 
 An acid is a substance containing hydrogen which may be replaced 
 wholly or in part by a metal. 
 
 A base is a substance containing hydroxyls which, ordinarily, may be 
 displaced wholly or in part by the acid radical. 
 
 In the preceding experiment we have seen that bases sometimes con- 
 duct themselves like acids. 
 
 Pb(OH) 2 + 2 Na OH Na-jPb O 2 + 2H 2 O. 
 
 As we proceed it will become evident that narrow rigid classifications 
 are not always possible. 
 
 Mono-basic, (HC1), Mon-acid, (NaOH), 
 
 Basicity, Di-basic, (H 2 SO 4 ), Acidity, Di-acid, Ca(OH) 2 , 
 
 Tri-basic, (H 3 PO 4 ). Tri-acid, Bi(OH) 3 . 
 
Il8 EXPERIMENTAL CHEMISTRY. 
 
 Oxacids ending in ic give ate salts. 
 
 Oxacids ending in ous give ite salts. 
 
 Oxacids with hypo and per affixes yield corresponding salts. 
 
 HC1O forms hypochlorites, 
 
 HC1O 2 forms chlorites, 
 
 HC1O 3 forms chlorates, 
 
 HC1O 4 forms perchlorates. 
 
 A metal may now be denned as a base-forming substance. NH 4 is 
 sometimes called a hypothetical metal. 
 
 A non-metal may be regarded as an acid-forming element. The 
 student should keep clearly in mind the relations between hydrogen and 
 acidic properties, and hydroxyls and basic properties. 
 
 PROBLEMS. 
 
 i. A liter of a given solution of NaOH contains 40.058 grm. of the 
 solute; an acid solution of unknown strength was titrated against the al- 
 kali solution when it was found that 50 cm. 3 of the acid solution was re- 
 quired to neutralize 25 cm. 3 of the alkali. What was the normality 
 (strength) of the acid solution? 
 
 2. How many grm. of H 2 SO 4 will be required to neutralize 1500 cm. 3 
 of a 2 N solution of KOH? To neutralize 400 cm. 3 of a 5 N solution of 
 Ba(OH) 2 ? 
 
 3. If 500 grm. of Na-jSO^ioELjO are prepared by neutralization, how 
 many grm. of acid were required ? Of alkali ? 
 
 4. If you were asked to determine the strength of an acid solution, 
 how would you proceed, assuming that you have access to a well equipped 
 laboratory ? 
 
 5. Define Alkalimetry. Acidimetry. Titer. Titration. 
 
CHAPTER XIII. 
 
 CHEMICAL EQUILIBRIUM, REVERSIBLE REACTIONS, MASS 
 LAW, DISSOCIATION. 
 
 "A chemical change in a given system is said to be complete when 
 it proceeds continuously with increase of 'products' and a correspond- 
 ing decrease of 'factors,' until one or more of the factors is exhausted 
 and the reaction ceases for want of more material. When increase in 
 quantity of products and decrease in the quantity of factors are arrested 
 before one or more of the factors is exhausted, it is known as an incom- 
 plete chemical change." Ostwald. 
 
 In the union of hydrogen and oxygen at various temperatures, to form 
 water, it will be recalled, that if the mixture of gases (2 vols. of H to i 
 vol. of O) was "sparked" at the ordinary temperature of the laboratory 
 the reaction continued with great speed until all the factors had been 
 exhausted, i.e., the reaction ran to an end. It will also be remembered 
 that if this quantity of water be heated to a temperature of 2500-3000 C., 
 it dissociates slightly into H and O gases. 2H 2 O * 2H 2 + O 2 . Now 
 it is further known, that in a closed system composed of H 2 O (steam) 
 at 3000 C. and H and O gases formed as the result of the decomposing 
 influence of heat, if the temperature is lowered the H and O tend to 
 recombine and form H 2 O, that is, the reaction runs to the left as shown 
 by the equation (i); if the temperature is again raised then equation (2) 
 represents the nature of the reaction. 
 
 (1) 2 H 2 O 2 H 2 + O. 
 
 (2) 2 H 2 O 2H, + O. 
 
 (3) 2 H 2 <=> 2H 2 + O. 
 
 It is evident that the H and O gases are in equilibrium with water 
 (steam), and the reaction is reversible. (Decompositions which are 
 reversible are called dissociations.) Increase of temperature tends to 
 produce a greater degree of dissociation; lowering the temperature 
 tends to produce a greater degree of association. What will be the 
 effect of increasing the pressure on the system ? Assuming that definite 
 volumes of H and O gases are in equilibrium with the steam, apply 
 pressure or introduce into the system volumes of either gas or both, then 
 equation (i) will represent the resulting reaction; if the pressure is 
 diminished or a portion of either gas, or both, is removed, equation (2) 
 shows the direction of the reaction. 
 
 A more detailed explanation may make the above phenomena more 
 easy of interpretation. Conceive of the steam at 3000 C. being placed 
 in a cylinder which has a movable piston and a stop cock. Keep the 
 
 no 
 
I2O EXPERIMENTAL CHEMISTRY. 
 
 temperature constant and diminish the pressure on the components by 
 pulling the piston well out toward the end, the steam will dissociate in 
 part into definite volumes of H and O always the same quantities under 
 identical conditions. Inside of the cylinder there is a definite con- 
 centration of each substance. The concentrations of the H and O gases 
 are very small, but not so small but that these particles meet one another 
 every now and then, and uniting, form water, but this immediately 
 diminishes the pressure on the remaining components of the system 
 for three volumes (2 of H and i of O) of hydrogen and oxygen gases 
 combine to form only two volumes of steam. Owing to this diminished 
 pressure more steam dissociates and the equilibrium is reestablished. 
 Now introduce through stop cock some H gas, this increases the pressure 
 upon the system, but especially does it increase the concentration of the 
 H particles, with the result that they meet the O particles more frequently, 
 and combination to form water is more rapid than under original con- 
 ditions. In other words, the speed of the reaction as shown by (i) is 
 much greater than the speed of the reaction indicated by (2), but after 
 the reaction (i) has continued for some time its speed will become less 
 owing to diminished concentration of the hydrogen particles. Oxygen 
 gas introduced, would have yielded similar results. Any factor which 
 alters the concentration of the "factors" or "products" will alter the 
 general direction of the reaction. If either of the components of the 
 system had been removed, the effect upon the nature of the ensuing 
 reaction would have been just as marked, only the reaction would have 
 proceeded as indicated by (2). 
 
 The above is one of many examples of concentration effect or effect oj 
 mass upon chemical reaction. So important is the influence of " active 
 mass" upon chemical reaction that it has been stated in the form of a 
 law by Guldberg and Waage. It is known as the Mass Law.* 
 
 " Interactions of substances depend not only upon the affinities in- 
 volved, but also upon the active mass oj the substances in a unit volume" 
 
 2H 2 O < 2 H 2 + O 2 . 
 C 2 = C 2 + C 3 . 
 
 I 2 
 
 C i 
 
 K is sometimes called the affinity constant. It really represents the 
 ratio of the affinities urging the opposed actions. The speed of the 
 
 reaction represented by (i) may be indicated by Speed! = Kj C i and 
 the opposed reaction by Speed 2 = K 2 C^.C 3 where K x represents the 
 
 * Professor T. W. Richards prefers to call it, " The Law of Concentration Effect. 
 
CHEMICAL EQUILIBRIUM, REVERSIBLE REACTIONS, ETC. 121 
 
 tendency for water to dissociate, and K 2 the tendency for H and O to 
 combine. 
 
 K, C 2 2 .C 3 C 2 2 .C 3 
 
 K.C'-K.C'.C,;- -;K = 
 
 K 2 C* C* 
 
 Experiment I. Reactions which are Approximately Complete. 
 
 (a) Formation of an inactive and practically insoluble product. Pour 
 3 cm. 3 or 4 cm.s of silver nitrate (AgNO 3 ) solution into a test tube, 
 (use side-shelf reagent); add a few drops of a dilute solution of sodium 
 chloride (NaCl); let the precipitate curd and settle to the bottom; observe 
 whether an increase in the quantity of the precipitant will cause additional 
 precipitation. Continue to add the precipitant until subsequent addi- 
 tions do not effect further precipitation of silver chloride (AgCl). Equa- 
 tion? Are the "products" of the reaction, active or inactive? Explain. 
 Is it possible that under certain conditions the "products" may react 
 to produce the "factors?" Are all substances soluble in degree? Is 
 the precipitate (AgCl) in equilibrium with the small quantity which is 
 dissolved ? 
 
 (b) Formation of a product (gaseous) which is removed. Prepare 
 a dilute solution of sodium carbonate (Na^CC^); test with litmus solution; 
 add a few drops of dilute sulphuric acid (H 2 SO 4 ) at a time and note effect 
 upon the speed of the reaction; continue to add acid until the effervescence 
 of CO 2 has ceased, and blue litmus begins to turn red. Equation? 
 When one of the products escape from the system, is it likely that the 
 "products" will react to produce the so-called "factors?" Is there any 
 likelihood of the establishment of an equilibrium between the "factors" 
 and the "products?" If not, is the reaction reversible? Explain. 
 
 Experiment II. Reactions which are Incomplete. Equilibrium. 
 
 (a) Reversible reaction. To 5 cm. 3 of a magnesium sulphate or 
 chloride solution (side-shelf), add a few drops of NH 4 OH at a time: 
 observe the partial precipitation of magnesium hydroxide (Mg(OH) 2 ); 
 now add an excess of NH 4 OH and note that subsequent additions of 
 NH 4 OH do not effect further precipitation. 
 
 (i) MgCl 2 + 2NH 4 OH <=*_M_g(OH)_ 2 + 2NH 4 C1. 
 
 Now add an excess of ammonium chloride, NH 4 C1 one of the "pro- 
 ducts," and observe that the presence of a large concentration of this 
 "active mass" causes reversal of reaction as indicated by (i). 
 
 Mg(OH) 2 + NH 4 C1 (excess) - MgCl 2 . 2 NH 4 Cl + NH 4 OH 
 MgCl 2 + NH 4 OH + NH 4 C1 (excess) -> MgCl 2 . 2 NH 4 Cl + NH 4 OH. 
 
 Inasmuch as Mg(OH) 2 is soluble in NH 4 C1 which is formed simul- 
 taneously, would you say that there is a complete precipitation in the 
 
122 EXPERIMENTAL CHEMISTRY. 
 
 first reaction? Do the "products'' tend to react and reproduce the 
 " factors?" When the speeds of these opposing reactions are the same, 
 a condition of equilibrium results. Was the speed of the reverse reaction 
 increased by increasing the concentration of one of the " active products " ? 
 If the NH 4 C1 is removed as rapidly as it is formed in first reaction, will 
 the reaction be complete ? Will there ensue a condition of equilibrium ? 
 Will the reaction be reversible? 
 
 (b) An incomplete reaction completed by rendering one of the active 
 products inactive by virtue of its insolubility in another medium. To 
 15 cm. 3 of a dilute solution of calcium chloride, CaCl 2 , add a few cm. 3 
 of a concentrated solution of potassium sulphate, K 2 SO 4 ; observe the 
 slight and partial precipitation of calcium sulphate, CaSO 4 . 
 
 CaCl 2 + K 2 SO 4 4= CaSO 4 + 2 K Cl. 
 
 Filter, and add 5 cm. 3 of ethyl alcohol, C 2 H 5 OH, to 5 cm. 3 of the 
 filtrate. Explain the formation of the heavy precipitate of CaSO 4 . 
 
 (c) A reaction involving equilibrium, in which each of the opposing 
 reactions are practically completed by varying the concentration of the 
 "active masses." 
 
 Add a few cm. 3 of an oxalic acid, H 2 C 2 O 4 , solution to 5 cm. 3 of a dilute 
 CaCl 2 solution; the precipitation of the calcium oxalate, CaC 2 O 4 , is 
 incomplete owing to the reaction of HC1 which is formed simultaneously. 
 
 CaCl 2 + H 2 C 2 O 4 <=>CaC 2 O 4 + 2 HC1. 
 
 Decant half of the solution upon a filter, and collect the filtrate. To 
 the portion remaining in the test tube add HC1 until precipitate dis- 
 solves. Write equation. To the filtrate add two or three cm. 3 of a 
 sodium acetate, NaC 2 H 3 O 2 , solution. Explain formation of precipitate 
 of CaC 2 O 4 . Write equations. Compare the reactions as represented 
 by equations. Are the reactions as represented by your equations, 
 opposite in nature? If these two opposing reactions should occur in 
 one system, with identical speeds, would a condition of equilibrium 
 results ? 
 
 Berthollet's Laws are in substance as follows: 
 
 " // the ' products ' are not active, or are active but are removed, then the 
 reaction is complete and runs to an end." 
 
 " When the products of a reaction are chemically active within a system 
 and are not removed, the reaction is reversible and incomplete, and results 
 ultimately in balanced action and chemical equilibrium." 
 
 Chemists are now inclined to regard all chemical changes as being 
 reversible, i.e., reactions do not run to completion. 
 
 A thorough understanding of Le Chatelier's Theorem will enable 
 the student to understand much in regard to chemical equilibrium which 
 would otherwise be vague. (See Energetics of Chemistry.) 
 
 Before leaving this subject the student is warned against looking 
 
CHEMICAL EQUILIBRIUM, REVERSIBLE REACTIONS, ETC. 123 
 
 upon a condition of chemical equilibrium as being due to a cessation or 
 suspension of chemical action. On the contrary such a state of balance 
 is due to the fact that the speeds of the opposing reactions are equal. 
 Again, a reversible action must not be considered as one which runs to 
 completion, when an opposite reaction is then initiated and runs back- 
 ward. Both actions are started at practically the same time; the one 
 gradually increases in speed and the other gradually diminishes in speed 
 until finally their speeds are identical, when a condition of equilibrium 
 results. 
 
 Experiment III. Dissociation of a Gas. 
 
 Place a small quantity of NH 4 C1 in a dry test tube; heat tube until 
 white fumes are evolved and sublime up the tube. Now hold a piece 
 of moistened red litmus paper in the mouth of the tube and it will be 
 turned blue showing that a portion of the ammonia is escaping; discon- 
 tinue heating tube for a moment and the blue litmus paper will be turned 
 red owing to the hydrogen chloride which is now escaping. Heat tube 
 strongly, and hold litmus paper in the heavy white fumes which appear 
 a few inches above the mouth of the tube. Results ? This is a case of 
 dissociation. Is temperature a factor in determining which way the 
 reaction shall go ? Explain. Would you say the heat oj dissociation is 
 positive or negative ? 
 
 NH 4 C1 <=> NH 3 + HCl 
 
 Many gases tend to dissociate when placed under favorable energy 
 conditions. At a temperature of 
 
 2 HI <= H 2 + I 2 
 N 2 O 4 <= NO 2 + NO 2 
 I 2 ^I + I 
 NH 4 C1 < NH 3 + HCl. 
 
 480 C. about 20 per cent, of the original quantity of HI is in the dis- 
 sociated condition. 
 
CHAPTER XIV. 
 NOTE ON THE MODERN THEORY OF SOLUTION. 
 
 The subject of "solutions" was introduced in a preceding chapter, 
 but little was said about the condition of the solute when in solution. 
 Attention was called to the fact that a substance in solution conducts 
 itself very much like a gas, and when the dissolved portion is in equilib- 
 rium with an undissolved portion, the solution is said to be saturated. 
 However, the mechanism of a solution was not considered in detail. 
 
 The purpose of these notes and the following experiments is to present 
 a comprehensive view of the modern theory of solution. 
 
 The theory has grown from the wrecks of other theories. Glimpses of 
 and approximations to the present theory logically preceded it. There- 
 fore, it is somewhat difficult to say just where it had its beginnings; how- 
 ever, it is certainly the result of bringing together, interpreting and corre- 
 lating a mass of apparently unrelated experimental data. 
 
 Without detracting one iota from the man who elaborated the theory 
 in practically its present form, we must acknowledge that we are indebted 
 to other men as well as to Arrhenius for the "Ionic Hypothesis" or The 
 Electrolytic Dissociation Theory. 
 
 Although the theory in its present form has been unable to meet in 
 an altogether satisfactory manner all of the requirements made of it, 
 yet it is quite possible when various relationships are more perfectly 
 understood that it will meet all demands. It is so vastly superior in its 
 ability to interpret an array of phenomena which would otherwise be 
 inexplicable that many noted investigators have called the theory "a 
 corner-stone of physical chemistry." 
 
 On investigating the labors of the chemists and physicists who pre- 
 ceded and were contemporaneous with Arrhenius, it appears as though 
 it would be perfectly justifiable to affirm that the theory as elaborated 
 was or could have been reached along practically two -independ- 
 ent lines of thought and investigation; namely, "electro-chemical" 
 and "osmotic and vapor pressure effects." That as it may or may not 
 be, the literature reveals Arrhenius, himself working along electro-chemi- 
 cal lines, as standing at the apex of two converging lines of investigation; 
 in fact, responsible for causing the lines to meet when they did. It was 
 Arrhenius who elaborated and enunciated the theory which correlated 
 a great volume of isolated data. Other men had been unable to do so. 
 It was Ostwald who then took up the theory, applied it, and finally be- 
 came its most effective and influential exponent. The theory now meets 
 with an almost universal acceptance. 
 
 To use the historical method of approach to our subject that is, to 
 
 124 
 
NOTE ON THE MODERN THEORY OF SOLUTION. 125 
 
 move historically along the two lines above enumerated would be per- 
 haps the most satisfactory, but this work is neither so voluminous nor 
 pretentious as to attempt more than to merely mention briefly the names 
 of the men who have contributed to the development of the theory. 
 
 Electro-Chemical. We shall not comment on the labors of Gilbert, 
 Dufay, Becaria, Priestly, Cavendish, Franklin, Wilke, Galvani, Volta and 
 Ritter, J. W., further -than to say that Volta originally separated all con- 
 ductors of electricity into two classes a first class, comprising the metals, 
 carbon and other good conducting substances found in nature, such as 
 metallic sulphides; the second class embraced such solutions as are con- 
 ductors of electricity. Until very recently we have spoken of the first class 
 of conductors as those in which when the current is flowing there is no 
 "simultaneous motion of the ponderable matter," while in the case of the 
 second class there is supposed to be a " corresponding motion of ponder- 
 able matter." That the flow of electricity through conductors of the first 
 class is convectional, as well as in the case of conductors of the second 
 class, is a recent theory that is supported by many. 
 
 About the year 1800 Nicholson and Carlisle found that when an electric 
 current was passed through two wires whose ends dipped into water, that 
 hydrogen and oxygen gases were liberated, and further, the liquid around 
 the pole at which the hydrogen was evolved was alkaline, while it became 
 acid around the other pole. 
 
 In 1802 Ermann secured results very similar to those of Nicholson and 
 Carlisle. Men attempted to explain the acid and alkaline reactions by 
 asserting that it was due to the effect of electricity upon water. 
 
 Sir Humphry Davy (1.778-1829) investigated the above phenomena 
 and demonstrated that pure water is separated into hydrogen and oxygen 
 gases, and that the acid and alkaline properties were due to impurities in 
 the water. Davy performed the experiment known to us as the "electroly- 
 sis of sodium sulphate," the solution having previously been colored blue 
 with a litmus solution. A satisfactory explanation of the phenomena was 
 not offered until many years later. Davy advanced that which may be 
 called the first "electro-chemical theory." It was based upon Dalton's 
 atomic hypothesis. In brief, his theory assumed that the atoms of the 
 elements when they came in contact with one another took to themselves 
 opposite charges of electricity; that if the charges were sufficiently strong 
 the two elements combined, forming a chemical compound. Also, if a 
 new and different atom came in contact with the compound and could 
 assume a sufficiently strong charge of electricity, it would attract the oppo- 
 sitely charged atom of the compound, causing decomposition. The 
 theory did not meet with a ready or even an ultimate acceptance. 
 
 Berzelius (1779-1848) now advanced his "dualistic theory" which in- 
 volved the "bipolarity of atoms." Affinity depended upon the size 
 and sign of the dominant electrical charge on the atom. Every com- 
 pound was conceived of as being composed of two parts, oppositely 
 charged the parts themselves might be composed of two parts. For 
 example, the salt, Zn SO 4 , should be written according to Berzelius, ZnO.- 
 
126 EXPERIMENTAL CHEMISTRY. 
 
 SO 3 , to show its real structure as being composed of a basic oxide and an 
 acid anhydride. (See Le Blanc's Electro-Chemistry.) 
 
 Grotthus (1805) announced the first complete theory of electrolysis, 
 known as the "exchange of partner theory." Clausius showed theory 
 to be a violation of the "second law of energetics." (Lehfeldt's Electro- 
 Chemistry contains a concise statement of the theory.) 
 
 Faraday (1794-1868) found that the quantity -of electricity passing 
 through a circuit and the chemical and magnetic effects produced were 
 proportional to one another. By a series of careful measurements Fara- 
 day, in 1833, established the first of two laws known as Faraday's Laws. In 
 substance the First Law is: the amounts of substances which separate at the 
 electrodes during electrolysis are strictly proportional to the quantity of 
 the electric current which passes through the electrolyte. The Second law 
 says: the mass of any substance liberated at the electrodes by a given 
 quantity oj electricity is directly proportional to the " chemical equivalent" 
 or combining weight. Davy believed that the current was carried through 
 solutions by little particles called "ions." He also injected into our nom- 
 enclature such terms, as electrolysis (the process of decomposing sub- 
 stances in solution by an electric current), electrolyte (the substance which 
 carries the current), anode (the positive electrode), cathode (the negative 
 electrode), ions (the particles into which the substance divides and act 
 as the real carriers of the current), anions (the ions which gather at the 
 anode), and the cations (ions which gather at the cathode). 
 
 Note. The student is reminded that although we shall retain these 
 terms, Faraday did not use them in the sense in which they are used at the 
 present time. 
 
 At that time there was considerable discussion as to what constituted 
 the anions and the cations. Berzelius said in the case of sodium sulphate, 
 which he wrote NaO and SO 3 , that NaO and SO 3 constituted respectively 
 the cations and anions that these reacted upon the water and produced an 
 alkali, NaO + H 2 O-> Na(OH) 2 , and an acid, SO 3 + H 2 O H 2 SO 4 , 
 but it was known that hydrogen and oxygen were liberated at the same 
 time. 
 
 Daniell disproved Berzelius's explanation by experimentally showing 
 that the hydrogen and oxygen liberated were chemically equivalent to the 
 alkali and acid formed. This would require, so to speak, " double electrical 
 action" which would be contrary to Faraday's laws. Daniell affirmed 
 that Na was the cation and SO 4 the anion; that these ions were set free at 
 the poles, and then reacted upon the water liberating hydrogen and oxy- 
 gen. " The salt alone must have conducted the electricity in the solution," 
 for if the water had conducted a part of the current of electricity there 
 would not have existed the interequivalence of acid and alkali, and 
 hydrogen and oxygen. 
 
 It became known later that both Ohm's law and Joule's law hold for 
 conductors of the second class as well as for the first class, therefore it 
 was obvious that none of the current spent in traversing an electrolyte 
 
NOTE ON THE MODERN THEORY OF SOLUTION. 127 
 
 was employed in doing chemical work in splitting up the electrolyte into 
 ions. Within the electrolyte during electrolysis the current merely exerts 
 a directive force on the ions; that is, so to speak, " sorts" them. 
 
 The theory of Grotthus, as mentioned previously, was found to be in 
 conflict with the " second law of energetics" by Clausius, who advanced 
 a theory of his own, in which he set forth the idea that the positive and 
 negative parts of a molecule in a solution are frequently in such a rapid 
 state of vibration that for a few moments at a time the parts may be said 
 to be independent of one another, so that if a current were passing at the 
 time the parts might be so favorably arranged as to follow the directive 
 force of the current and be deposited on electrodes. In other words, 
 he assumed a constant interchange of the negative and positive parts of 
 the molecule, and at any instant there was a small number of these parts 
 of the molecules momentarily free, and that the current was carried by 
 these molecule-parts. This theory was commonly accepted and is the 
 one which preceded our present theory. 
 
 Hittorf began his work about this time, on what he called the " migration 
 of the ions." He found during the electrolysis of a CuSO 4 solution be- 
 tween electrodes of copper, that the solution became very much more con- 
 centrated around the anode then around the cathode, yet the quantity of 
 copper deposited at the cathode was greater than the amount lost by 
 the cathode chamber. The conclusion was not difficult that the ions of 
 Cu and SO 4 had different migration velocities. From the speed ratios 
 Hittorf calculated the "transport numbers." 
 
 Kohlrausch, as the result of his work on the conductivity of dilute solu- 
 tions, established a very simple relation between the " transport numbers " 
 and the molecular conductivity. Kohlrausch's work also emphasized the 
 fact that inorganic substances, such as acids, bases and salts, were good 
 conductors (electrolytes), and organic substances were either non-conduc- 
 tors or very poor conductors. 
 
 As the result of the labors of Hittorf and Kohlrausch, Svante Arrhenius 
 (1887), a Swedish chemist, was enabled to reach certain conclusions re- 
 garding the theory of "free ions." Arrhenius while working on the con- 
 ductivity of solutions recognized two kinds of "molecules" one active 
 and the other inactive as he named them because of the belief that only 
 the "active" molecule was instrumental in causing conductivity. He 
 stated further his belief that the "inactive" changed into "active" 
 molecules with sufficient dilution. He was not able to offer a convincing 
 proof until he was placed in receipt of certain data from the other line 
 of investigation referred to before, namely, " osmotic and vapor pressure " 
 effects. 
 
 It had not been so very long before this time that investigators had ac- 
 cepted with considerable hesitancy "the dissociation of molecules" as 
 the explanation of the deviations of many vapor densities from that antici- 
 pated by theory. Planck at this time as the result of purely thermody- 
 namical considerations said that "it would be quite natural to expect a 
 similar dissociation when substances were placed in solution." 
 
128 EXPERIMENTAL CHEMISTRY. 
 
 Osmotic Pressure. Without attempting to make a complete review of 
 all the work that had been done along this line up to the time of Arrhen- 
 ius, it will be dismissed with the statement that the question of solution 
 had been holding the attention of eminent investigators for a period of 
 ten to fifteen years preceding the time at which Arrhenius announced his 
 theory. The fundamental idea which stimulated men to these investi- 
 gations was that when a substance is in the dissolved condition it behaves 
 in a manner very similar to a gas. Although the idea was not new, yet it 
 secured the attention of Van't Hoff (now in the University of Berlin), 
 who became firmly convinced of the fact that a " substance in solution 
 behaves like a gas and, like it, exerts a pressure" but he was unable to 
 devise a method for the measurement of the pressure. 
 
 We are told that Van't Hoff was one day walking with De Vries, the 
 zoologist, and remarked about the pressure which forces upward the sap 
 in a tree, whereupon De Vries replied that Pfeffer, the botanist, had in- 
 vestigated and also measured the force. This was the very data for 
 which Van't Hoff had been searching. 
 
 Pfeffer had observed some years previously that when a plant cell was 
 placed in very dilute solutions it bursted, and when placed in concen- 
 trated solutions, the cell shriveled. At the time Pfeffer was unable to 
 make any satisfactory measurements, but later, due to the labors of Traube, 
 he was able to deposit semipermeable membranes of copper-ferrocyanide 
 in the walls of porous porcelain cups which he connected with a mano- 
 meter; the cup was then filled with a sugar solution and immersed in 
 pure water. As the result of Pfeffer's experiments, although he was una- 
 ware of their far-reaching influence, it was found that the pressure (osmo- 
 tic) was directly proportional to the concentration and the absolute tem- 
 perature of the solution. Van't Hoff was quick to correlate osmotic 
 pressure and gas pressure. He found, in other words, that the gas laws 
 hold approximately for osmotic pressure. Further, in substance, that 
 equal fractions of the molecular-gram weights of organic substances dis- 
 solved in the same volume of water produce the same osmotic pressure. 
 To state his conclusions from another point of view and include Avo- 
 gadro's rule, "At the same osmotic pressure and temperature equal 
 volumes of all solutions contain the same number of molecules, and, in 
 fact, that number which under the same pressure and at the same temperature 
 exists in the same volume of gas." As Van't Hoff proceeded with his work 
 he found that strong inorganic acids and bases and salts did not act nor- 
 mally, i.e., the osmotic pressures as yielded by solutions of these substances 
 were sometimes more than double that which he anticipated; for example, 
 when a given weight of hydrogen chloride was dissolved in water it gave 
 an osmotic pressure equal to almost twice that which it should yield ac- 
 cording to the gas law (PV = RT). He was unable to account for this. 
 
 Freezing Point and Vapor Pressure of Solvents. About this time 
 Raoult made two important generalizations in regard to dilute solutions 
 which came as the result of the labors of himself, Bab, Wullner, Ostwald, 
 Blagden, Rudolph and others. Our knowledge of the depression of the 
 
NOTE ON THE MODERN THEORY OF SOLUTION. I2Q 
 
 freezing point and elevation of the boiling points of solvents by dissolved 
 substances, was uncertain and fragmentary before Raoult initiated his 
 investigations. (See chapter on "Solutions.") 
 
 General Law of the Vapor Pressure of Solvents. i molecule oj a non- 
 saline, non-volatile substance, dissolved in 100 molecules of any volatile 
 liquid, lowers the vapor pressure oj this liquid by a nearly constant fraction 
 oj its value approximately .0105. 
 
 The experience of Van't Hoff relative to osmotic pressures was the 
 experience of Raoult in regard to deviations from his law concerning the 
 freezing point, namely, that organic compounds gave normal and com- 
 paratively uniform results, while strong inorganic acids and bases and 
 salts gave a larger depression than was anticipated. Raoult was like- 
 wise unable to explain it. 
 
 In the year 1887 Van't Hoff contributed an article to the Zeitschrift fiir 
 physikalische Chemie on " The Role of Osmotic Pressure in the Analogy 
 Between Solutions and Gases." The article was in support of his be- 
 lief that the gas law was applicable to dilute solutions of substances. 
 It called attention to the work of Rauolt and emphasized the fact that 
 the use of strong acids and bases caused deviations from both his own 
 and Raoult's laws. Near the latter end of the article Van't Hoff says, 
 " * * * Arrhenius pointed out to me, by letter, the probability that salts 
 and analogous substances when in solution break down into ions. As a 
 matter of fact, as far as investigation has been carried, the solutions which 
 obey the law of Avogadro are non-conductors, which indicates that they 
 are not broken down into ions; and a further experimental examination 
 of other solutions is possible, since, from the assumption made by Arrhen- 
 ius, the deviation from Avogadro's law can be calculated from the con- 
 ductivity." 
 
 As soon as the above article appeared, Arrhenius was able, by comparing 
 the effects of electrolytes and non-electrolytes in depression of the freezing 
 point of water with their respective electrical conductivities, to produce 
 convincing proof of his original assumption of the electrolytic disso- 
 ciation of certain salts in aqueous solution. The organic compounds of 
 Van't Hoff and Raoult were the non-electrolytes of Kohlrausch and 
 Arrhenius. 
 
 Arrhenius set forth his proof and theory in an article entitled, " Uber 
 die Dissociation der im Wasser gelosten Stoffe." He called attention to 
 the fact that the compounds, salts, strong acids and strong bases which 
 create abnormally high osmotic pressures and abnormal depressions of 
 the freezing point are all electrolytes i. e., when dissolved in water 
 their solutions become good conductors of electricity; on the other hand, 
 the substances which give normal osmotic pressures and normal depres- 
 sions of the freezing points are either non-electrolytes or conduct the 
 current very poorly. A normal solution of cane sugar, which is a non- 
 electrolyte, freezes at 1.87 C.; a normal solution of sodium chloride 
 which is a good electrolyte freezes at 3.46 C. A very similar relation 
 holds in regard to the osmotic pressures of the two solutions. It is gener- 
 9 
 
130 EXPERIMENTAL CHEMISTRY. 
 
 ally accepted that the pressure of a gas or the osmotic pressure or the de- 
 pression of the freezing point is directly proportional to the number of 
 molecules or ultimate particles present in the gas or solution. Arrhenius 
 concluded, therefore, that if electrolytes produced abnormal alterations in 
 these particular properties, that the molecules of the substance in solution 
 split up (dissociated) into a larger number of particles which, so far as 
 above phenomena are concerned, acted like molecules. Each of these 
 particles would produce its own effect on osmotic pressure or depression of 
 the freezing point and that inasmuch as those substances which behaved 
 normally with reference to osmotic pressure, etc., were non-electrolytes, 
 these smaller particles were the real cause of the conductivity of a solution. 
 Arrhenius gave the general name ions to these independent particles into 
 which molecules were assumed to dissociate. The terminology of the 
 theory is very similar to that suggested by Faraday; in fact, many parts 
 are identical. 
 
 Theory. In briefest outline, his theory is as follows: Whenever an elec- 
 trolyte, referring to the solute, is dissolved it almost wholly or partly disso- 
 ciates into ions. The extent of the dissociation depends upon nature of so- 
 ute, concentration, solvent, temperature and pressure. These ions many be 
 composed of one or more atoms, but are altogether different in nature from 
 the elementary substances of which they are composed. These ions 
 are electrically charged, whereas the elementary substances are neutral, 
 therefore their energy content is different, and logically we expect a differ- 
 ence in properties. An electrolyte dissociates into two ions one (cation), 
 bearing a positive, and the other (anion), a negative charge of electricity. 
 Frequently one of the ions further dissociates or it may associate with 
 another ion forming a complex ion. It should be remembered that the 
 solution as a whole is electrically neutral and therefore, there must be as 
 many electrical charges of one kind as there are of another. Irrespective 
 of the masses of the ions, the number of charges carried by each ion is 
 equal to the valency of the atom or atomic group which constitutes the 
 ion. As a matter of fact, on the electrical theory, valency really amounts 
 to nothing more or less than the "number of positive or negative unit 
 charges associated with the chemical atom." An example may make the 
 theory more easily understood. When hydrogen chloride is dissolved in 
 water it immediately dissociates into ions of hydrogen (hydrion) and chlor- 
 ine (chloridion), respectively known as cations and anions. The ions are 
 wholly independent of one another. Each ion retains its identity, i.e., 
 properties of reaction, velocity, etc., regardless of its birth. Dilution in- 
 creases the extent of the dissociation. If the solution is examined, there 
 will be no evidence of any free atoms of either hydrogen or chlorine gas. 
 If these gases were in the free state they would be very easy to detect as 
 hydrogen is very slightly soluble and chlorine is a yellow colored gas. 
 The fact that no electrical energy is consumed in dissociating the mole- 
 cule within the liquid, and yet when a current is passed through the solu- 
 tion the hydrogen ions (cations) are directed toward, and will appear, at 
 the negative electrode (cathode), while the chlorine ion (anion) is attracted 
 
NOTE ON THE MODERN THEORY OF SOLUTION. 131 
 
 to the positive electrode (anode), is strong evidence that the electrolyte 
 immediately dissociates on being placed in solution and that the ions 
 are the actual carriers of the current. When the ions reach their respective 
 electrodes their charges of electricity are unloaded and they are trans- 
 formed into the elementary substances; they are then deposited upon the 
 electrode, escape as a gas or react chemically upon the water. Since 
 hydrogen has a valency of one it is said to have one electrical charge 
 and that positive; chlorine 
 
 HC1, aq. <= H + Cl' NaOH, aq. + Na + OH' 
 
 ions have one negative charge. 
 
 H 2 SO 4 , aq. < H + HSO' 4 NaCl, aq.<= Na + Cl' 
 
 lt + 
 S0" 4 + H . 
 
 It is evident that in an electrolytic solution the solute is partly in the molec- 
 ular condition and partly in the ionic state. Water is not a good con- 
 ductor of electricity, yet it is very slightly dissociated. 
 
 (99-95%)H 2 0^H + + OH'(.o 4 i%) 
 
 The remark is frequently made, however, that if it is slightly acidified with 
 sulphuric acid it becomes a good conductor. Interpreted in terms of 
 the "electrolytic dissociation" theory, this means merely that the acid is 
 dissociated into hydrogen and the acid radical ions which carry the 
 current. 
 
 The student should not confuse the ideas of solubility and electrolytic 
 dissociation. Great solubility does not imply great dissociation. 
 
 For purposes of convenience all substances have been divided into three 
 groups electrolytes, half -electrolytes and non-electrolytes. The term 
 "electrolyte" has been frequently used in two different senses some- 
 times referring to the solute and at other times used to designate the con- 
 ducting solution. It is more properly used in this latter sense. The 
 word " iongen " has been proposed as a suitable name for those substances 
 which dissociate when placed in solution. This seems to be more consis- 
 tent with facts, as salts, acid and bases, with the exception of fused salts, 
 are in themselves non-conductors of the electric current. Solutions of 
 salts and strong acids and bases are included in the class known as elec- 
 trolytes; solutions of such weak acids and bases, as acetic acid and ammo- 
 nium hydroxide, which are poor conductors of the electric current, are 
 grouped as half-electrolytes, while all other substances, largely organic, 
 which do not conduct are called non-electrolytes. 
 
 Before proceeding with special phases and applications of the theory, 
 it might be well to consider a number of explanations which have been 
 volunteered in an effort to answer the question as to the source of the elec- 
 tric charges possessed by the ions. 
 
 Source of Electric Charges. Le Blanc thinks that in a compound like 
 
132 EXPERIMENTAL CHEMISTRY. 
 
 potassium iodide there remains some chemical energy after the metallic 
 potassium and iodine have chemically combined, and that this chemical 
 energy, through the influence of the solvent, is transformed into electrical 
 energy which is seated in the charges of the ions. By the aid of an 
 electric current during electrolysis it is possible to add to these ions the 
 energy which they originally possessed as neutral substances, when they 
 will separate in the ordinary molecular forms at the electrodes. 
 
 According to Nernst, the dielectric constant of a liquid and its dissociat- 
 ing power are in direct proportion. He holds that the forces which bind 
 the ions together to form a molecule are due to electrical attractions be- 
 tween the oppositely charged ions. The specific inductive capacity of 
 water at 20 C. is 81.1 times that of air; therefore, since the ions are sup- 
 posed to have a constant charge, the force exerted between the ions in a 
 molecule, tending to prevent the dissociation will be 81.1 times less in 
 water than in air. The dielectric constant of ethyl alcohol is 26, for 
 ethyl ether 4.36, and for carbon disulphide 2.6. 
 
 The "electron theory," as advanced by J. J. Thomson, suggests"that 
 there is only one kind of ultimate particle the electron, an atom of elec- 
 tricity which has a mass af about one-thousandth that of a hydrogen atom. 
 This theory necessitates a revision of our former notions of the chemical 
 atom. It assumes "that atoms of the various elements are collections 
 (constellations) of these electrons positive and negative charges held 
 together mainly by their electric attractions, associated with more or less 
 ether." These corpuscles are conceived of as being in a state of rapid 
 vibration. In some atoms the velocities of the electrons (negative charges) 
 may be so great that a corpuscle escapes at once from the atom, thus 
 leaving the atom positively charged; or the reverse may be the case, leaving 
 the atom negatively charged. A negatively charged body contains an 
 excess of electrons; a positively charged one, a deject. It is thought, for 
 example, that when sodium and chlorine, each representing a particular 
 "constellation" of electrons, combine to form sodium chloride that the 
 two groups of electrons are modified. Now, it is difficult to break up a 
 molecule into atoms out of which it was formed, but comparatively easy 
 to split it in a slightly different manner, so as to leave one electron in 
 excess with the chlorine, and a deject of one with the sodium, with the re- 
 sult that they become charged positively and negatively, "respectively. 
 The organic chemist is familiar with the fact that it may be quite difficult 
 to split a long carbon compound at a particular bond or link, but very 
 easy perhaps to break it at some other link. It follows from the above 
 that "since electrons possess a definite though very small mass, the asso- 
 ciation with, or removal of, the electric charges from matter causes a differ- 
 ence in its mass. 
 
 d + e _ Q' 
 
 Na e Na 
 
 + 
 
 Chlorine has an atomic mass of 35.45; therefore, the mass of a chlorine 
 ion would be 35.451, and the weight of a sodium ion would be .001 less 
 
NOTE ON THE MODERN THEORY OF SOLUTION. 133 
 
 than the atomic weight. Here as elsewhere the student should not con- 
 fuse fact and theory. 
 
 In an article by G. T. Beilby, which recently appeared in The Chemi- 
 cal News, London, some very interesting suggestions are made by the 
 author in regard to the probable source of the energy of solution. The 
 following are extracts from said article: " The solute molecules in a dilute 
 solution of any non-volatile solid, are solid molecules sparsly distributed 
 among a multitude of intensely active solvent molecules, the temperature 
 of the solution being frequently many hundred degrees below that at 
 which they could of themselves assume the greater freedom of the liquid 
 or gaseous state. These solute molecules have to a great extent been set 
 free from the constraining effect of their cohesive forces, but it is important 
 to remember that this freedom has not been attained by the increase of their 
 own kinetic energy as in liquefaction by heat. Their freedom and the 
 extra kinetic energy they have acquired have in some way been imparted 
 to them by the more active solvent molecules; for, if the solvent could be 
 suddenly removed, leaving the solute molecules still similarly distributed 
 in a vacuous space, they would eventually condense into a solid aggregate. 
 This must be the case, for the non-volatile solute has no measurable vapor 
 pressure at the temperature of the solution. The kinetic energy of the 
 solute molecules is of itself quite insufficient to endow them with the prop- 
 erties of the gaseous or even of the liquid molecule, even when their 
 cohesive forces have been weakened or overcome by separation. 
 
 // the energy employed in this separation is not intrinsic to the solute 
 molecule then it must in some way have been imparted by the solvent mole- 
 cules. It therefore becomes important to compare the energy endowment 
 of one set of molecules with that of the other." * * * "Taking into con- 
 sideration not only this greater store of energy, but also the much smaller 
 cohesive force of water as compared with the majority of solid solutes, 
 there can be no doubt that the active role in aqueous solutions must be 
 assigned to the solvent, not to the solute molecules. 
 
 This leads to the important conclusion that the energy oj solution, oj 
 diffusion, and of osmosis is due, not to imaginary gaseous energy oj the 
 solute, but to the actual liquid energy oj the solvent" * * * "The rude 
 mechanical jostling to which the complex molecule is subjected will natu- 
 rally tend to break it up into simpler portions which are mechanically 
 more stable." 
 
 "The view that the phenomena of solution depend on the relative 
 kinetic energy of the solvent and solute molecules appears to apply with 
 special force to the phenomena of dissociation in dilute solutions. Under 
 the gas theory there does not appear to be any reason why the solute 
 molecules should dissociate into their ions. So obvious is this absence of 
 any physical motive that Professor Armstrong has happily referred to 
 the dissociation as " the suicide of the molecules." " With the acceptance 
 of the view that phenomena of solution are largely due to kinetic energy 
 of the solvent molecules, the phenomena of dissociation also appear to 
 take their place as a natural result of this activity, for consider the situa- 
 
134 EXPERIMENTAL CHEMISTRY. 
 
 tion of an isolated molecule (solute), closely surrounded by and at the 
 mercy of some millions of water molecules all in a state of intense activity." 
 
 "An ideally perfect solution that is, a solution of which the physical 
 properties are determined solely by the number of molecules it contains 
 in a given volume must consist of a solvent and a solute which have no 
 chemical affinity for each other, so that their molecules will neither asso- 
 ciate not dissociate in solution. Probably comparatively few solutions 
 will be found which even approximate to this ideal perfection." 
 
 It appears as though the two latter theories might be correlated so as 
 to become supplementary. 
 
 References. 
 
 Theoretical Chemistry, Nernst. Inorganic Chemistry, Ostwald. 
 
 Physical Chemistry, Walker. Lehrbuch d. Allg. Chem., Ostwald. 
 
 Physical Chemistry, Jones. Scientific Foundations, Ostwald. 
 
 Electro-Chemistry, Lehfeldt. Scientific Memoirs, containing orig- 
 
 Electro-Chemistry, Le Blanc. inal papers of Pfeffer, Van't Hoff, 
 
 Theory of Electrolytic Dissocia- Raoult, Arrhenius. Edited by Ames. 
 tion, H. C. Jones. 
 
 APPLICATION OF THEORY. 
 Electrolysis and Electrical Equivalents. 
 Experiment I. Electrolysis. 
 
 (a) Refer to experiment in which the electrolysis of H 2 O was performed 
 through the agency of Na 2 SO 4 . Interpret the phenomena observed in 
 the light of the "ion theory." Represent the electrolytic reactions by 
 means of equations in which are shown the ions. Explain the formation 
 of free elements at the electrodes. 
 
 (b) Perform the electrolysis of HC1. Explain mechanism. 
 
 Careful measurements have shown that 96,580* coulombs (one "Fara- 
 day") of electricity will deposit one gram-equivalent of any substance. 
 This quantity of electricity is known as the electro-chemical constant. A 
 short rule for determining the gram-equivalent is: divide the atomic 
 weight of an element, or the sum of the atomic weights if it is a radical, 
 by its valency; the quotient will represent the number of grams of that 
 substance deposited by 96,580 coulombs. The electro-chemical equivalent 
 may be found by dividing the gram-equivalent by 96,580. It is the amount 
 deposited by one coulomb. Example: Atomic weight of oxygen is 16; 
 its valency is 2; therefore 8 grams is the gram-equivalent. For silver it 
 is, 107.93 "*" I = IO 7-93 grams. The electro-chemical equivalent of silver 
 is, 107.92 -v- 96.580 = .0011175 grams. Recall Faraday's laws. 
 * Richards', T. W., value. 
 
NOTE ON THE MODERN THEORY OF SOLUTION. 135 
 
 Names of Ions. The following list is based on Walker's system: 
 
 Anions. Cations. 
 
 Cl' Chloridion Na' Natrion (Sodium) 
 
 CIO/ Chloranion K' Kalion 
 
 CIO/ Perchloranion NH/ Ammonion 
 
 NO/ Nitranion Ca" Calcion 
 
 SO/' Sulphanion H' Hydrion 
 
 SO 3 " Sulphosion Fe'" Triferrion 
 
 S" Sulphidion Fe" Diferrion 
 
 HC 2 H 3 O/ Acetanion Cu" Dicuprion 
 
 OH' Hydroxidion Ag' Argention 
 
 Some of the most interesting things which are learned in regard to ions 
 by means of experiments in electrolysis are, (a) no electrical energy is 
 consumed in the production of ions; their birth is antecendent to the pass- 
 ing of the current, (b) they carry electrical charges which are proportional 
 to their relative valencies, (c) their relative migration velocities differ 
 greatly, but the migration velocity of a particular ion is independent of the 
 nature of the co-existent ions, and (d) the absolute velocity of the most 
 speedy ion is very slow. The following absolute velocities at 18 calcu- 
 lated for infinitely diluted aqueous solutions are given by Kohlrausch. 
 The difference of potential between the electrodes i cm. apart was i volt. 
 
 K = .00066 cm. H = .00320 cm. 
 
 NH 4 = .00066 cm. Cl = .00069 cm. 
 
 Na = .00045 cm. NO 3 = .00064 cm. 
 
 Li = .00036 cm. C1O 3 = ,00057 cm - 
 
 Ag = .00057 cm - OH = .00181 cm. 
 
 It is seen that the movement of the ions through practically pure water 
 is very slow. Walker calculates the force" required to drive i gram of 
 hydrion through water at the rate of i cm. per second to be equal to about 
 320,000 tons weight. It is obvious that the ions experience great resist- 
 ance to their movements. This is said to be due, at least in part, to the 
 " hydration of the ion," i.e., the comparatively large amount of water which 
 travels with the ions. 
 
 CONDUCTIVITIES. 
 
 Experiment II. (L. T.) Electrolytes, Half -Electrolytes and Non-Elec- 
 trolytes. 
 
 (Instructions.) By means of a number of conductivity cells (Figs. 
 26 and 27) and an alternating current suitably applied, determine the 
 relative electrical conductivities of the following substances. On basis 
 of results arrange substances under the three heads suggested above. 
 Sugar (C 12 H 22 O U ) solution, toluene (C 7 H g ), NaCl solution, C 2 H 5 OH, 
 chloroform (CHC1 3 ), HC1 solution, NaOH solution, acetic acid (HC 2 H 3 O 2 ) 
 solution, NH 4 OH, distilled H 2 O, dilute solution of C 2 H 5 OH, NH 4 C1 solu- 
 tion, HNO 3 solution, KC 2 H 3 O 2 solution, C 6 H 6 . 
 
i 3 6 
 
 EXPERIMENTAL CHEMISTRY. 
 
 Note. It is suggested that .iN or .5N solutions of acids, bases and salts 
 be used in above experiment. 
 
 In view of your experiments, what would you say of the conductivity 
 of organic compounds as compared with inorganic substances in aqueous 
 solution ? 
 
 There is, accurately speaking, no sharp line of demarcation between these 
 classes. The distinction is based, however, on degree of conductivity. 
 It should be recalled that electrolytes produced abnormal osmotic press- 
 
 FIG. 26. 
 
 FIG. 27. 
 
 ures, abnormal depressions of the freezing point and abnormal lower- 
 ings of the vapor pressure; that is, they (solutes) apparently dissociated in 
 water in a manner similar to the dissociation of gases under favorable 
 conditions of pressure and temperature. 
 
 N 2 4 <=> N0 2 + N0 2 
 
 On the other hand, the non-electrolytes yield normal results relative to 
 osmotic pressure, depression of freezing point, and lowering of vapor 
 pressure. It seems, therefore, that conductivity is due to these little 
 particles into which the molecules are dissociated, namely, anions and 
 cations. And because these ions are attracted to the positive and nega- 
 tive electrodes respectively of an electrolytic cell, they must be oppositely 
 charged. This kind of splitting up of the molecule has been called 
 "electrolytic dissociation" to distinguish it from gaseous dissociation 
 where the particles are not apparently electrically charged. 
 
 Other conditions being the same, conductivity depends upon the degree 
 of ionization, i.e., the number of ions present. 
 
 Experiment III. Ionization in Solution. 
 
 (a) Nature of Solute. Determine relative electrical conductivity 
 of aqueous solutions of the following: NaCl, C 2 H 5 OH, HNO 3 , CHC1 3 , 
 
NOTE ON THE MODERN THEORY OF SOLUTION. 137 
 
 KOH, NH 4 OH, KNO 3 , NaOH, Na,,SO 4 , Ca(OH) 2 , HC1, HC 2 H 3 O 2 and 
 HgCl 2 . Tabulate data under the following heads: acids, bases, salts, 
 organic compounds, so that substances .having greatest conductivities 
 will stand at the head of their respective columns. 
 
 Is it evident from above that the ionization of different solutes vary 
 greatly ? Write the formula for one substance in each column showing 
 the ions formed by its dissociation. What ions are common to all acids ? 
 Bases? If the "strength" of acids and bases depends upon the relative 
 concentrations of hydrogen and hydroxidions respectively, which of the 
 above represent the " strong " acids and bases ? Indicate by a star placed 
 opposite the formula. 
 
 As a general rule, it may be stated that aqueous solutions of salts and 
 the so-called "strong" acids and bases are good conductors because they 
 are largely ionized. Mercuric chloride is an exception, but not the only 
 one. Dissociation of a substance may be said to be one of its specific 
 properties, and, as such, it would be expected that different substances 
 possess this property in varying degrees. The beginner is again cautioned 
 against concluding that because a substance is very soluble that it is very 
 largely dissociated. Sugar is quite soluble, but we have no evidence that 
 it dissociates. 
 
 (b) (i) Effect and nature of solvent. Proceed as in foregoing ex- 
 periments to ascertain the relative conductivity of: (i) distilled H 2 O, (2) 
 dry NaCl, (3) dry C 12 H 22 O n , (4) dry Nal, (5) aqueous solutions of (2), 
 (3), (4). Using C 2 H 5 OH, CHC1 3 , and C 7 H 8 instead of water repeat (5). 
 
 Add alcohol to a- NaCl solution from time to time, and as frequently note 
 effect on the conductivity of the solution. Results ? 
 
 (b) (2) Prepare some perfectly dry HC1 gas by bubbling it through 
 concentrated H 2 SO 4 ; run the gas into toluene (C 7 H 8 ) or benzene (C 6 H 6 ) 
 until latter is saturated, then run gas into distilled H 2 O; test each of the 
 solutions with litmus paper. Results ? Now place the solutions of HC1 
 in two separate and thoroughly cleaned and dried conductivity cells and 
 determine relative conductivity. Results ? Does the toluene solution of 
 HC1 show any acid properties whatever ? Are there any ions in this solu- 
 tion ? Your reasons ? Set aside this latter solution for future use. 
 
 What would you say of the dissociating power of H 2 O as compared 
 with other solvents tested ? 
 
 The nature of the solvent plays a very important role in determining 
 whether a resulting solution will conduct or not, i.e., whether the solute 
 will or will not ionize. Therefore conductivity does not depend alone on 
 either solute or solvent. Water is the most efficient dissociating agent 
 with which we are acquainted. 
 
 At ordinary temperatures pure substances appear to have small con- 
 ductivities, providing the transfer is electrolytic. At higher temperatures 
 fused salts and metallic oxides are good electrolytic conductors. It was 
 maintained at one time that only aqueous solutions were good electrolytes, 
 but it has since been learned that certain non-aqueous solutions frequently 
 give higher conductivities. Liquefied NH 3 , SO 2 and HCN dissolve 
 
138 EXPERIMENTAL CHEMISTRY. 
 
 salts and make good conductors. Their properties, however, are much 
 more complex than those of water. 
 
 A satisfactory and final explanation as to why water or any other solvent 
 acts in such a manner has not been volunteered. It has been suggested 
 as a general rule: 
 
 " When the solvent is associated, the dissolved molecules are mostly 
 simple, and vice versa." 
 
 There is experimental evidence to support the idea that in pure water 
 there is found in addition to H 2 O <= H* + OH', polymerized or asso- 
 ciated molecules of water with formulae probably varying from H 4 O 2 
 H 8 O 4 . Water itself when pure conducts only in a minimum degree. 
 It should be kept in mind that the hydrions and hydroxidions of dissoci- 
 ated water may unite with other oppositely charged ions, and the com- 
 pound thus formed may or may not dissociate depending upon its nature 
 and the solvent. 
 
 It has been calculated that ten to eleven million liters of pure water 
 contain one gram of hydrions and seventeen grams of hydroxidions. 
 
 When NaCl is dissolved in water a portion of it dissociates as 
 per the equation, NaCl <= Na' + Cl', but owing to 
 
 H 2 O = O H' + H- 
 
 JI Jf 
 NaOH HC1, 
 
 the dissociation of water the above equations represent the reactions that 
 will undoubtedly occur, as well as the number of different particles present. 
 
 (c) Effect of dilution. Test the conductivity of glacial acetic acid or 
 concentrated sulphuric acid. Add a few drops of water; test. Repeat 
 operations many times. Record all results. Is ionization increased or 
 diminished by dilution? Write equations showing the dissociation on 
 addition of water to HC 2 H 3 O 2 . How many particles are present in the 
 solution? Write similar equations for NaOH and NajSO^ 
 
 It is quite obvious that if increased conductivity is due to an increase in 
 ionization in the above experiment that dilution produces a greater degree 
 of dissociation. 
 
 Arrhenius established the following law which although not absolutely 
 correct is sufficiently accurate for practical purposes. 
 
 " The degree of dissociation of a substance in a solution is equal to 
 the ratio of the equivalent conductivity of that solution to its equivalent 
 conductivity at infinite dilution." 
 
 A v 
 
 Another law of much importance to the chemist is known as Ostwald's 
 Law of Dilution. It applys to a binary electrolyte, i.e., when each mole- 
 cule forms two ions. The law involves the mass law. 
 
NOTE ON THE MODERN THEORY OF SOLUTION. 139 
 
 " The product of the concentration of the two ions divided by the con- 
 centration of the undissociated part is a constant" 
 This law may be represented in formula as follows: 
 
 (1) Na' + Cl'^ 
 
 (2) C 3 . C 2 = K 
 
 C 2 .C 3 
 
 (3) K 
 
 KG, 
 
 is known as the "dissociation constant" and its value varies with each 
 solute and solvent and temperature. It states, however, a constant re- 
 lation which exists between the ions and the undissociated molecules of 
 an electrolyte. The relation of dilution to degree of ionization may be 
 shown by altering the form of the equation. Let y equal the fraction 
 that is dissociated, then i y equals the part which is in the molecular 
 condition, and v the volume to which the solution is diluted. 
 
 y y i y 
 
 (4) C s = - ; C 2 = - ; Q = -- . 
 
 V V V 
 
 y 2 
 
 y 2 
 
 (s) ' - y 
 
 = K; (i y) v = K. 
 
 It is evident on mere inspection that if we dilute the electrolyte, i.e., 
 increase its volume (v) by addition of the pure solvent, then we diminish 
 all the concentrations of the substances represented in the formula 
 see (3); but this will affect the numerator more than the denominator 
 because the former is a product of two concentrations. If such were true 
 then K would not remain constant, which is contrary to our experimental 
 data. In order then for the value of K to remain constant the concentra- 
 tions of C 2 and C 3 must become greater through the dissociation of the 
 undissociated molecules represented by Q. The formula therefore rep- 
 resents that dilution produces a greater degree of ionization. 
 
 Although the law of dissociation as given by Ostwald does not apply to 
 strong electrolytes, yet the formula may be taken as practically correct up 
 to a concentration of .oiN. 
 
 The dissociation constant K is a characteristic of every compound. 
 If the degree of ionization is large for any substance then the value of 
 K will be large and vice versa. We shall see later how this "constant" 
 is really a measure of the strength of acids. 
 
 The form of the law of dilution suggests a condition of equilibrium as 
 
140 EXPERIMENTAL CHEMISTRY. 
 
 existing between the product of the ions on one side and the molecules on 
 the other. As a matter of fact this is the real statement of the relations 
 existent. The ions are in equilibrium with the undissociated molecules. 
 If the concentration of any of the particles is altered, say C 2 is increased, 
 then some of C 3 will combine with C 2 to increase the value of C x until the 
 mathematical form of the equation is re-established. Again, if the value of 
 C x should be increased, the reaction would run in such a direction as to 
 increase the values of C 2 and C 3 until original conditions prevail. Such 
 changes are called reversible and the change itself is known as a balanced 
 action. 
 
 The effect of temperature and pressure upon ionization is easily antici- 
 pated by applying Le Chatelier's Theorem. In some cases the heat of 
 dissociation is positive while in others it is negative. 
 
 COLOR OF IONS. 
 
 The theory of electrolytic dissociation requires that the color of an elec- 
 trolyte, i.e., a dilute salt solution shall depend upon the color of its free 
 ions. There is much experimental evidence to support this view. 
 
 Experiment IV. Persistency of the Color of an Ion. 
 
 (a) Observe the colors of solutions of the following substances as 
 you find them among the " shelf reagents." Write the equations showing 
 the ions formed by dissociation. Tabulate results, stating the color of 
 the respective anions and cations: 
 
 NaCl, NaBr, KNO 3 , KC1, Na 2 SO 4 , KC 2 H 3 O 2 , NH 4 C1, KI, BaCl 2 , 
 CaCl 2 , KBr, NH 4 C 2 H 3 O 2 . 
 
 (b) Now examine solutions of Cu(C 2 H 3 O 2 ) 2 , CuCl 2 , Cu(NO 3 ) 2 , CuSO 4 ; 
 compare the color of the solutions with the substance in the dry solid 
 state; write equations showing ions formed by dissociation; by means of 
 (a) tabulate colors of anions; what ion is the cation common to all the 
 solutions ? What is the probable source of the color common to the solu- 
 tions? Is the color likely due to the salt in the molecular condition? 
 Why ? What is the color of anhydrous copper sulphate ? (Recall a for- 
 mer experiment.) Note colors of solutions of Co(NO 3 ) 2 , CoCl 2 , K 2 CrO 4 . 
 State source of color and your reasons for thinking so. 
 
 (c) Place a small pinch of CuBr 2 in a test tube. What is the color 
 of the dry salt ? Add a few drops of water. Observe color. Add water 
 to test tube until one-fourth full. Is it probable that the dominant color 
 is due to molecular CuBr 2 ? Your reasons? Now fill the test tube 
 with water. What is the color ? Is it similar in color to solutions of cop- 
 per salts used in (b) ? Why was it necessary to dilute the solution to so 
 great a degree to procure final color? Write the formula for Ostwald's 
 Law of Dilution. Interpret above experiment in terms of it. 
 
 (d) Effect of magnitude of electrical charge on color of ion. Prepare 
 dilute solutions of FeSO 4 , and FeCl 3 . Show by means of formulae the ions 
 
NOTE ON THE MODERN THEORY OF SOLUTION. 141 
 
 formed in the respective solutions. What is the color of the anions in 
 both? Is the iron ion common to both solutions? Is its electrical 
 charge (valence) identical in both cases ? Is its energy content identical 
 in both solutions ? Would you infer that the latter had something to do 
 with its properties? What are your conclusions in regard to the differ- 
 ence in color of the two solutions ? Are the two iron ions identical ? In view 
 of (a), (b), (c), would you consider it an extravagant assumption to assume 
 that any particular ion always maintains its identity, i.e., possesses an in- 
 dividual set of chemical and physical properties? 
 
 Ostwald in particular has studied this question of the absorption spectra 
 of solutions from the stand-point of the dissociation theory, and the fol- 
 lowing conclusion was reached: 
 
 "That the spectra of dilute solutions of salts containing the same 
 colored ion are identical." 
 
 CHEMICAL CONDUCT OF IONS. 
 
 With comparatively few exceptions chemical reactions are the result 
 of the mutual interaction of ions. 
 
 Experiment V. Ions Necessary to Chemical Reaction Acids. 
 
 (a) Place a portion of the concentrated H 2 SO 4 or HC 2 H 3 O 2 , which 
 would not conduct the electric current, upon a piece of dry zinc. Results ? 
 Dilute the acids with four or five times their volume of distilled water. 
 Results? Explain and write equations showing action of water upon 
 acids and of solutions upon zinc. 
 
 (b) Pour some of the toluene saturated with dry HC1 upon a dry piece 
 of marble. Results? Did the above solution act like an electrolyte? 
 Add a little water to a portion of the toluene solution of HC1; test its 
 conductivity; test its action upon a dry piece of marble. Results? Ex- 
 plain. Equations ? 
 
 (c) Recall or repeat the experiment of mixing dry NaHCO 3 and 
 H 2 C 4 H 4 O 6 in a mortar. Observe whether any chemical action takes place. 
 Add water to the mixture. Results ? Explain in terms of the dissocia- 
 tion theory. Equations? 
 
 Concentrated acids are usually very slightly ionized; that is, they yield 
 but few ions, their conductivity is low and their chemical activity as an 
 acid is correspondingly weak. Acids are substances which when dis- 
 solved in water or other dissociating solvents, yield hydrions. This hydrion 
 is common to all acids and is the source of the acidic properties of the 
 general class of substances known as acids. The hydrion is a colorless 
 substance composed of one atom of hydrogen bearing one electrical 
 charge. It has a sour, acid taste and turns blue litmus red. It may 
 transfer its charge to some metals like zinc and magnesium and suffer 
 displacement from the solution. Zn + 2H' + SO" 4 Zn" +-SO" 4 + 
 
142 EXPERIMENTAL CHEMISTRY. 
 
 Hj. It combines with the hydroxidion to form water. Its migration ve- 
 locity is much greater than that of any of the other common ions, there- 
 fore it confers large conductivities upon those solutions in which it is con- 
 tained, providing there is sufficient concentration of said ions. 
 
 It follows that the "strength" of an acid will depend upon the con- 
 centration of hydrions which it can supply. In other words, those acids 
 whose "dissociation constants" (K) are large, are known as "strong" 
 acids and good conductors, and vice versa. Conductivity measurements 
 are approximate measurements of the relative "strength" of acids. 
 
 Experiment VI. "Strength" of Acids. 
 
 Determine the conductivities of the following acids in the concentrated 
 and in the diluted conditions: H 3 PO 4 , H 2 SO 4 , HC1, HC 2 H 3 O 2 , HNO 3 . 
 Tabulate results, arranging acids in order of increasing conductivity. 
 
 Any process which removes hydrions from a solution or prevents their 
 appearance in same will eliminate all acid properties. Toluene prevented 
 the dissociation of HC1, hence neither conductivity nor acid reaction of 
 the toluene solution of hydrogen chloride. 
 
 Experiment VII. " Strength " Of Bases. 
 
 Proceed as in former experiments to ascertain the relative conductivities 
 of iN. solutions of NH 4 OH, NaOH, Ca(OH) 2 , and KOH. Tabulate re- 
 sults. Write an equation for each substance showing nature of dissociation. 
 Which ion is common to all the solutions ? 
 
 As in the case of acids, we find solutions of bases giving varying con- 
 ductivities, depending largely upon their relative degree of ionization. 
 Bases are substances, the aqueous solutions of which yield hydroxidions. It 
 is this ion which confers the basic properties upon the class of substances 
 known as bases. The hydroxidion is a colorless substance composed of 
 one atom each of hydrogen and oxygen (hydroxyl), bearing one electrical 
 charge. It possesses a caustic, lye-like taste and turns red litmus blue. 
 It combines with the hydrion to form water. Its migration velocity, 
 though much greater than other ions, is but little more than one-half the 
 velocity of the hydrogen ion. The "strength" of a base will depend 
 upon its "dissociation constant." "Strong" bases will have a constant 
 whose value is comparatively large; "weak" bases, the reverse. 
 
 HC1 - H- + Cl' 
 HN0 3 IT + N0' 3 
 HBr H- + Br' 
 NaOH OH' + Na ! 
 KOH OH' + K- 
 LiOH OH' + Li- 
 
NOTE ON THE MODERN THEORY OF SOLUTION. 143 
 
 Neutralization. Salts. 
 
 H- + Cl' + Na- + OH' Na- + Cl' + H 2 O + 13700 cals. 
 H- + Cl' + K- + OH' K- + Cl' + H 2 O + 13700 cals. 
 H- + Cl' + Li' + OH' Li' + Cl' + H 2 O + 13700 cals. 
 
 H- + OH' H 2 O + 13700 cals. 
 
 H- + Cl' + NH- 4 + OH' NH- 4 + Cl' + H 2 O + 12200 cals. 
 H- + C 2 H 3 0' 2 + Na- + OH' Na- + C 2 H 3 O 2 + H 2 O + 13300 cals. 
 H- + F' + Na- + OH' -> Na' + F' + H 2 O + 16200 cals. 
 
 Neutralization is the process whereby the hydrions of an acid and 
 hydroxidions of a base combine to form water. 
 
 THERMO-CHEMICAL SUPPORT OF THE DISSOCIATION THEORY. 
 
 The "heat of neutralization" of strong acids and bases furnishes one 
 of the strongest supports for the "ion theory." Jones says: 
 
 Since all processes of neutralization of completely dissociated acids and 
 bases are the same, the heat of neutralization of all such acids and bases 
 must be a constant, and must be the heat of combination of a gram equiv- 
 alent of hydroxyl and hydrogen ions. 
 
 It may be added, if either the acid or base is "weak," the heat of 
 neutralization will not be 13700 cals, but will vary from that value; it may 
 be more or less. 
 
 Change in Volume Support. Another bit of evidence said to support the 
 theory of solution is the uniform contraction in volume observed during the 
 neutralization of strong acids and bases. When 1000 cm. 3 of a normal solu- 
 tion of a strong base are used to neutralize 1000 cm. 3 of a strong acid, the 
 resulting mixture has a volume which is always 20 cm. 3 less than the sum 
 of the two original volumes. The inference usually made is the i gram 
 of hydrions and 1 7 grams of hydroxidions occupy a larger volume by 20 cm. 3 
 than 1 8 grams of water which they form as the result of the neutralization. 
 
 A salt is formed during the process of neutralization, and if the latter 
 is complete it may be said to be the result of the union of the cation of 
 the base and the anion of the acid. 
 
 Salts are named, as a rule, according to composition and without re- 
 gard to their conduct. Recall former classification. 
 
 Experiment VIII. Salts. 
 
 (a) Formation of an insoluble salt. To a few cm. 3 of AgNO 3 solution 
 procured from the "shelf reagents," add 5 cm. 3 of distilled water. Shake 
 vigorously. Divide into four equal portions, using test tubes as recep- 
 tacles. To one portion add a few drops of HC1, to another a little CHC1 3 , 
 to another a few drops of a NaCl solution, to the remaining tube add a 
 small quantity of a KC1O 3 solution. Record all data. Give equations 
 showing the interaction of ions. 
 
 (b) Formation of a complex salt. To a few cm. 3 of AgNO 3 solution 
 add two or three drops of KCN. Caution. KCN is a poison. Add a 
 
144 EXPERIMENTAL CHEMISTRY. 
 
 little more KCN; continue to add until a clear solution is obtained. Add 
 a few drops of HC1. Result ? Explain various reactions. Give equations. 
 (c) Repeat above using a solution of CuSO 4 ' Continue to add KCN 
 until solution is colorless. Explain. Equations? 
 
 A "complex salt" is one which gives a complex ion, for example, 
 KAg(CN) 3 K- + Ag(CN)' 2 . A "compound salt" like K 2 Mg(SO 4 ) 2 
 manifests less tendency to form complex ions. K 2 Mg(SO 4 ) 2 * 2K* 
 + Mg" + 2SO' 4 . It would be in accordance with facts to state that 
 both "complex" and "compound" salts go through successive steps of 
 ionization. 
 
 HYDROLYSIS. 
 
 Heretofore we have spoken of water as a neutral substance, but this is 
 far from being true. It will be recalled that w r ater dissociates into hy- 
 drions and hydroxidions. If a salt when dissolved in water simply under- 
 goes dissociation the solution is neutral, since there are not present any 
 appreciable amounts of free hydrions and hydroxidions. If, however, the 
 nature of the cation of the salt is such that it forms a weak base when 
 combined with an hydroxyl group, it will combine with the OH ions of the 
 water, and being a weak base, it remains to a great extent in the associated 
 condition. If it should be only slightly soluble, it will be precipitated. 
 This removal of cation of salt and OH ions of water, leaves the anion 
 of the salt and the H ions of the water in the solution, which now possesses 
 an acid reaction due to the free hydrions. If the anion of the salt forms 
 a weak acid when combined with a hydrion, then by a similar line of 
 reasoning this would leave free hydroxidions in solution which would then 
 possess an alkaline reaction. 
 
 Experiment IX. Hydrolysis. 
 
 (a) Place 5 cm. 3 of distilled water in a test tube; test it with litmus 
 paper to assure yourself that it is neutral. Dissolve a small crystal of 
 A1C1 3 or A1 2 (SO 4 ) 3 . Are these "neutral salts"? Test the solution. 
 Results ? Explain. Equations ? 
 
 (b) Using distilled water and N^COg repeat above experiment. Re- 
 sults ? Explain. Equations ? 
 
 (c) Place a small crystal of SbCl 3 in the bottom of a test tube; add 
 5 cm. 3 of distilled water; shake vigorously, then add concentrated HC1 
 drop by drop until solution becomes clear after warming. Add from 
 25 to 50 cm. 3 of water. Results? Explain. Give equations. 
 
 (d) Turn to the experiment in your note book under "Salts," in which 
 you tested with litmus paper solutions of various salts. Some of the 
 results were rather perplexing at the time; can you explain away your 
 previous difficulties? 
 
 Hydrolysis may be denned as a case of double decomposition in which 
 water is one of the chemical reagents. 
 
NOTE ON THE MODERN THEORY OF SOLUTION. 
 
 Per Cent Ionized. 
 (Calculated at 18 C.) 
 
 Acid. 
 
 Per Cent. 
 
 Acid. 
 
 Per Cent. 
 
 HC1 cone 
 
 1 3 60 
 
 HI 5N 25 C. 
 
 oo oo 
 
 HC1 iN 
 
 . . 78 . 40 
 
 H 2 F 2 , iN 
 
 . 7 .00 
 
 HNO 3 , cone. . . . 
 HNO 3 , iN 
 H 2 SO 4 cone. 
 
 9 . 60 
 82.10 
 o 70 
 
 H HS, .iN 
 H.H 2 P0 4 , . 5 N 25 
 HC 2 H 3 O 2 , iN. . . 
 
 06 
 C....I7.00 
 .40 
 
 H 2 SO 4 iN 
 
 ' 
 
 51 20 
 
 HC 2 H 3 O 2 .iN 
 
 i 40 
 
 H.HC 2 O 4 , .iN 25 
 H HCO 3 iN 
 
 c 50.30 
 
 O 12 
 
 HC 4 H 4 6 , iN 25 
 HCN, .iN . . 
 
 C.... 8.20 
 
 O.OI 
 
 
 
 
 
 Base. 
 
 Per cent. 
 
 Base. 
 
 Per Cent. 
 
 Li OH, iN 
 NaOH, iN 
 
 63 . oo 
 . .73.80 
 
 Ca(OH) 2 , N/6 4 , 
 Sr(OH) 2 N/64, 2 
 
 25 C. 90.00 
 
 < C. Q3 .OO 
 
 KOH, iN. 
 
 . . 78 .00 
 
 Ba(OH) 7 , .iN,. 
 
 . . 7 <; . oo 
 
 K OH .iN 
 
 86.00 
 
 Ba(OH) 9 , N/64, 
 
 2Z C. 02 .OO 
 
 NH OH iN 
 
 o 40 
 
 NH 4 OH .iN 
 
 I . 4O 
 
 
 
 
 
 Salts. 
 
 Per Cent. 
 
 Salts. 
 
 Per Cent. 
 
 NaCl iN 
 
 67 . 60 
 
 CuSO 4 iN 
 
 22 .OO 
 
 NaCl, ,5N 
 NaCl iN 
 
 -73-40 
 83 oo 
 
 NH 4 C1, iN 
 
 Zn SO 4 iN 
 
 74.00 
 .24.00 
 
 NaHCO 3 , iN... 
 NaC 2 H 3 O 2 , iN.. 
 
 52-00 
 53 .00 
 
 Zn C1 2 , iN 
 Hg C1 2 , iN 
 
 48 . oo 
 0.90 
 
 Na 2 S0 4 , iN 
 KC1 iN 
 
 44-50 
 75 oo 
 
 Hg(CN) 2 , iN... 
 Ag NO,, iN.. 
 
 Small 
 . .58.00 
 
 KNO 3 , iN 
 
 64 . oo 
 
 Ca SO 4 , .oiN . . 
 
 63.00 
 
 KC 2 H 3 O 2 iN 
 
 64 oo 
 
 BaCl 2 iN 
 
 75 .00 
 
 K 2 SO 4 , iN 
 
 2 OO 
 
 MgSO., .iN.. 
 
 45 .00 
 
 K 2 CO 3 iN 
 
 40 OO 
 
 KC1O 3 -5N 
 
 . . 70 .00 
 
 
 
 
 
 * Water, o.ooooi per cent, ionized. 
 
 IO 
 
146 EXPERIMENTAL CHEMISTRY. 
 
 Experiment X. Ionic Equilibrium. 
 
 Prepare a saturated solution of NaCL Add concentrated HC1 slowly 
 drop by drop, until a white granular precipitate of NaCl appears. Write 
 formulae showing nature of the dissociation of the NaCl and the HC1. 
 Is there an equilibrium between the ions and the undissociated molecules 
 of a solute? Allow the salt to settle; decant the clear supernatant liquid 
 into another test tube; add C 2 H 5 OH until a fine granular precipitate of 
 NaCl appears. Explain. 
 
 To interpret correctly the above phenomena it is necessary to recall 
 certain principles presented previously. An example may facilitate the in- 
 terpretation. When NaCl is dissolved in water the ions are in equilibrium 
 with the molecular salt, and further, at a definite temperature there is a 
 constant relation between the two as indicated by the equations which 
 follow: 
 
 NaCl < Na- + Cl' HC1 <=> H" + Cl' 
 
 L-j L, 2 C 3 v_x 4 C 5 C 6 . 
 
 KQ = C 2 .C 3 KC 4 = C 5 .C 6 . 
 
 This is a mathematical expression of the equilibrium. If the solution 
 is saturated then the right and left-hand members of the equation pos- 
 sess their maximum values, for K is constant, i.e., is independent of the 
 concentration of the solution. If the solution is diluted then the value of 
 Q becomes smaller. Hence it follows, that in a saturated solution at a 
 given temperature for a given solute the concentration of the undisso- 
 ciated molecules, and the product of the concentrations of the ions are 
 constant. KQ is called the "solubility product" and is a characteristic 
 of each substance. C 2 x C 3 is sometimes spoken of as the "concentration 
 product" or "ion product." If the " solubility product" is exceeded, 
 supersaturation or precipitation will result. 
 
 It might be well to mention in this connection that a precipitate is in 
 equilibrium with the undissociated molecules remaining in solution. 
 
 Na' + Cl'+ NaCl (Dslvd.) 
 
 u 
 
 NaCl. (Ppt) 
 
 Referring again to the above experiment, the value of C 3 was greatly 
 increased by C 6 , with the result that C 2 , in order to preserve the mathe- 
 matical form of the equation, became correspondingly small by associ- 
 ating with some of C 3 and C 6 to form NaCl. This in turn increased the 
 value of Q. When this latter exceeded the " solubility product " of NaCl, 
 salt was precipitated. 
 
 SOLUTION TENSION. 
 
 Experiment XI. Displacement of Ions by a "Free Metal." 
 
 (a) Dissolve 5 grams of Pb(C 2 H 3 O 2 ) 2 in 100 cm.3 of distilled H 2 O; 
 place solution in a small Erlenmeyer flask; suspend a piece of zinc in 
 
NOTE ON THE MODERN THEORY OF SOLUTION. 147 
 
 the clear solution; set flask aside to stand for several hours. Note the 
 displacement of lead by the zinc. 
 
 Pb" + 2C 2 H 3 O' 2 + Zn->Pb + Zn" + 2C 2 H 3 O / 2 . 
 
 (b) Repeat (a) using CuSO 4 solution instead of Pb (C 2 H 3 O 2 ) 2 . Results? 
 Equation ? 
 
 (c) Repeat (b) using a piece of clean iron wire bent into the form of a 
 coil. Results ? Equation ? 
 
 (d) Dissolve a little HgNO 3 in hot water, then dip a coiled piece of 
 copper wire into the solution. Set aside for an hour. Results? Rub 
 wire with piece of filter paper. Results ? Equation ? 
 
 (e) Pour 10 cm. 3 of AgNO 3 solution into a crystallizing dish; add a few 
 drops of mercury; set aside for an hour or more. Results? Equation? 
 
 In all of the above experiments were the displaced ions, anions' or 
 cations ? 
 
 Experiment XII. Displacement of Ions by a "Free Non-Metal." 
 
 (a) To separate solutions of KI and KBr, or Nal and NaBr, add a 
 few drops of "chlorine water." Results? Add a little CHC1 3 to each 
 test tube and shake. Results? Explain. Equations? 
 
 (b) To a solution of KI add a little "bromine water." Results? 
 Add a little CHC1 3 and shake. Results ? Explain. Equations ? In 
 the above experiments were the displaced ions, cations or anions? 
 
 The inferences which are drawn as the result of the foregoing experi- 
 ments are, (a) the metals manifest varying tendencies to press into solution, 
 i.e., the ionic condition, since they are not otherwise soluble; (b) the non- 
 metals relative to one another, behave in a manner very similar to the 
 metals. It is possible to arrange the metals in a series in the order of 
 their "decreasing solution tension." 
 
 Alkali metals Lead 
 
 Alkaline earth metals Hydrogen 
 
 Magnesium (Arsenic) 
 
 Aluminum Copper 
 
 Manganese , Antimony 
 
 Zinc Bismuth 
 
 Chromium Mercury 
 
 Cadmium Silver 
 
 Iron Palladium 
 
 Cobalt Platinum 
 
 Nickel Gold 
 Tin 
 
 As a general rule, the ions of any metal are displaced from a normal 
 solution by any of the free metals which precede it in the above series. 
 
148 EXPERIMENTAL CHEMISTRY. 
 
 The question undoubtedly arises as to why metals have any such 
 property as "solution tension." The following experiments and discus- 
 sion may enable the student to understand a little more clearly the nature 
 of this phenomena. 
 
 Experiment XIII. Chemical Energy and Electrical Energy. 
 
 (a) Prepare a cold dilute aqueous solution of H 2 SO 4 (4 of H 2 O to i of 
 acid). Pour the solution into a small beaker. Suspend a strip of clean 
 zinc in the solution; by means of a copper wire connect the zinc with one 
 of the binding posts of a galvanometer or a volt-meter. Read volt-meter. 
 Is there evidence of much chemical action between zinc and H 2 SO 4 ? 
 Bind another copper wire to which is attached a copper strip to the other 
 binding post; suspend copper strip in the solution in such a manner as not 
 to touch zinc strip. Read volt-meter. Results? Is there more or 
 less evidence of chemical action in the beaker than when previously 
 examined? Observe and record phenomena manifested by this typical 
 cell. Did either the solution or the metals contain electricity originally ? 
 What is the probable source of the electrical energy? Place a small 
 quantity of "granulated" zinc in the beaker containing the acid solution 
 after removing the strips of Cu. and Zn. ; test the escaping gas by means of 
 a lighted match. Results? When chemical action has ceased, remove 
 excess of zinc and once more insert the strips of Zn. and Cu. ; take the 
 reading of the volt-meter. Explain. The beaker now contains a solu- 
 tion of what? Evaporate a portion of the solution to dryness. Is it 
 CuSO 4 ? Your reasons? 
 
 (b) The above experiment may be repeated by substituting for zinc 
 the following metals, iron, copper, lead, aluminum and others as 
 desired. Arrange the metals in order of their increasing electro-motive 
 force. (E.M.F.). 
 
 The student is familiar with the fact that when heat energy is contributed 
 to solids or liquids in sufficient quantities, the molecules pass into the 
 space above the liquid, and if it be a closed system, equilibrium is estab- 
 lished for a given temperature, when the pressure of the vapor is equal 
 to the vapor-tension of the substance. Increase the amount of internal 
 energy of the substance, i.e., raise its temperature, and it is obvious that its 
 vapor-tension is also increased. 
 
 Every metal has a certain solution-tension which tends to push its 
 particles into solution in the form of ions. The metal will continue to 
 dissolve until its solution-tension is in equilibrium with the solution-press- 
 ure of the ions. Since energy can not come from nothing, it is said that 
 the chemical energy or the free energy of the system is transformed into 
 the electrical charges of the ions. Metals yield cations, and non-metals 
 anions. 
 
 Now it is maintained by Nernst and others that the solution-tension of 
 substances is analogous to vapor-tension. An application of the theory 
 
NOTE ON THE MODERN THEORY OF SOLUTION. 149 
 
 of "electrolytic solution-tension" will reveal the points of analogy as well 
 as the differences. 
 
 If we dip a piece of metal, say zinc, into pure water, then owing to the 
 solution-tension of the metal some ions will pass into solution. These 
 ions are charged with positive electricity. Since both kinds of electricity 
 must be simultaneously produced whenever electrical energy comes into 
 existence, the metallic zinc becomes negatively electrified. The solution 
 is now positively electrified and the metal negatively electrified. The 
 positively-charged ions are attracted by the negatively-charged metal, so 
 that a difference of potential is established. The metal will, however, 
 continue to press into solution, thereby acquiring a lower potential, until, 
 as mentioned above, the solution-tension is in equilibrium with the solu- 
 tion pressure of the ions. The solution and the metal now have their 
 greatest difference of potential. The "single potential" of metals im- 
 mersed in pure water obviously depends upon their respective solution- 
 tensions. Suppose we now dip a metal like silver into a solution of its 
 ions which have a greater ionic solution-pressure than the metal's solution- 
 tension, this will result in the deposition upon the metal of some of its ions. 
 The metal will become positively electrified and the solution negatively 
 electrified. 
 
 Again, if we place a metal in a solution whose ionic solution-pressure 
 is just equal to its solution-tension, equilibrium is established at once. 
 Ions will neither be formed nor deposited, hence there will be no differ- 
 ence of potential between metal and solution. 
 
 In the above discussion we have referred alone to the solution-tension 
 of metals which form positive ions. Le Blanc says, that as far as we know, 
 all substances capable of yielding negative ions have a high solution- 
 tension. This would leave a free non-metal like chlorine positively 
 charged. 
 
 We immediately infer that as the result of solution-tension, metals and 
 non-metals acquire a definite electro-motive force, when placed in pure 
 water or a solution (usually a normal solution) containing their respective 
 ions. This E.M.F. is frequently called the " single potential" of metals 
 or non-metals. 
 
 " The tendency toward chemical reaction and its accompanying trans- 
 formation of energy in a cell is measured by the E. M. F. or voltage. 1 ' 
 Lehfeldt's Electro- Chemistry. 
 
 We have considered the questions of solution-tension, the production of 
 "single potentials" and the E.M.F. of a cell, but we have not considered 
 in detail the probable source of energy of which the foregoing are mani- 
 fested forms. This energy is evidently, at least in part, derived from the 
 chemical energy of the system. It was thought for some time that the 
 whole of the chemical energy was converted into electrical energy, but 
 this assumption is by no means justified by the facts. This is only true 
 in the case of cells whose E.M.F. do not vary with temperature. In cer- 
 tain types of cells a portion of the chemical energy is evolved as heat, the 
 cell becoming warmer as the cell continues in action; in other types the 
 
150 EXPERIMENTAL CHEMISTRY. 
 
 electrical work done exceeds the chemical energy spent, i.e., heat is ab- 
 sorbed and transformed into electrical energy. 
 
 " Measurement of electro-motive force may then be looked upon as a 
 means of determining the change of the free energy in a chemical reac- 
 tion." Lehfeldt's Electro-Chemistry. 
 
 The E.M.F. is a measure of the chemical energy of a system when the 
 process is reversible, i.e., when the system is neither warmed nor cooled 
 by the reaction which produced the E.M.F. 
 
 TABLE OF SINGLE POTENTIALS. 
 
 (The free metal is placed in a normal solution of its ions.) 
 
 K = (+ 2.9) H=- .277 
 
 Na = (+ 2.54) Cu = - .606 
 
 Ba = (+ 2.4) As = - .62 ? 
 
 Sr - (+ 2.3) Bi =- .67 ? 
 
 Ca = + 2.28 Sb = - .74 ? 
 
 Mg = + i .21 Hg = - i .027 
 
 Al = + i . oo Ag = - - i . 048 
 
 Mn = + .80 Pd = - 1.07 ? 
 
 Zn = + .493 Pt = - 1.14 ? 
 
 Cd = + .143 Au = - - 1.35 ? 
 
 Fe = + .063 I = + .797 
 
 Tl = + .045 Br = + 1.270 
 
 Co = . 045 O = + i . 396 ? 
 
 Ni = - .049 OH = + 1.396? 
 
 Sn = - .07 ? Cl = + 1.694 
 
 Pb = - .129 NO 3 == +1.75 
 SO 4 .= +1.90 
 
 It will be observed that the relative positions of the jree metals in the 
 " solution- tension" series and the " E.M.F." series are identical. The 
 order in which the metals occur in this series is especially significant, 
 inasmuch as it represents the relative chemical activities of the jree metals. 
 All of the metals in the fore part of the list readily oxidize when exposed to 
 the air; while those in the latter part of the list do not. It follows from 
 this that the former metals will be reduced with difficulty from their ox- 
 ides, while the reverse will be true of the latter. The arrangement ex- 
 presses the combining relations of the metals with regard to other elements 
 as well as oxygen. None of the metals in the fore part of the list are 
 found free in nature. It should be remembered in this connection that all 
 of these same metals can displace hydrogen from an acid, while those 
 which succeed hydrogen in the list and are unable to displace hydrogen 
 are found free in nature. Other relationships will be found represented 
 by this series. 
 
NOTE ON THE MODERN THEORY OF SOLUTION. 
 
 HEAT OF IONIZATION. 
 
 Lithium = + 261000 joules. 
 Potassium = + 257000 joules. 
 Sodium = + 237000 joules. 
 Strontium = + 244000 joules. 
 Calcium = + 226000 joules. 
 Magnesium = + 225000 joules 
 Aluminum = + 165000 joules. 
 Manganese = + 102000 joules. 
 Zinc = + 69500 joules. 
 Cadium = + 34600 joules. 
 Iron (ferrous) == + 42400 joules. 
 Thallium = = -f 3400 joules. 
 Cobalt = + 31700 joules. 
 
 Nickel = + 29400 joules. 
 Tin == + 2900 joules. 
 Lead = + 2900 joules. 
 Hydrogen = 2300 joules. 
 Copper = 37200 joules. 
 Mercury (Hg') = 86900 joules. 
 Silver = 110900 joules. 
 
 Chlorine = + 167400 joules. 
 Bromine = + 121000 joules. 
 Iodine = + 57500 joules. 
 Oxygen (1/4 O 2 ) = + 88600 joules. 
 
 Ostwald (Lehrbuch. d. allg. Chem.) has calculated the "heat of ioni- 
 zation" of one equivalent for a number of metals. The above table is 
 quoted from his work. 
 
 " The heat of ionization of hydrogen is so small as almost to lie within 
 the margin of errors, and may be ignored. Hence, the approximate rule, 
 that the heat of ionization of a metal is practically equal to its heat of 
 solution in dilute (i.e., completely dissociated) acid." Lehfeldt's Electro- 
 chemistry. 
 
 After inspecting above table it is evident that some of the ions possess 
 more available energy than when in the condition of the free metal; other 
 ions, the reverse. 
 
CHAPTER XV. 
 THE NON-METALS, OR ACID-FORMING ELEMENTS. 
 
 THE HALOGEN GROUP. 
 
 (THE CHLORINE FAMILY.) 
 Fluorine, F, 19 
 Chlorine, Cl, 35.45 
 Bromine, Br, 79.96 
 Iodine, I, 126.97 
 
 These four elements compose what is known as a "natural group"; 
 the members are closely connected by a similarity of chemical proper- 
 ties; i.e., they resemble one another in their chemical relations, and 
 by combination with identical substances, produce a series of corre- 
 sponding compounds which resemble each other in their respective 
 chemical properties. 
 
 The members of this group are usually termed the halogens, because 
 of their tendency to produce salts resembling sea-salt in their composition. 
 These salts are called haloid salts or simply, halides. 
 
 FLUORINE, F. 
 At. Wt. 19 Mol. Wt. 38. 
 
 Preparation and Properties. This element is not found free in nature. 
 The chief sources are calcium fluoride, CaF 2 , commonly known as fluor- 
 spar and cryolite, A1F 3 3NaF. The latter mineral is found in relative 
 abundance in Greenland. 
 
 Fluorine is a pale yellowish-green gas with an unpleasant odor like 
 chlorine. Although fluorine does not combine with oxygen it shows a 
 great affinity for hydrogen, with which it unites explosively, even in the 
 dark. In fact, so great is the affinity of fluorine for hydrogen that many 
 compounds which contain the latter are decomposed when brought into 
 contact with it. It is recorded that Moissan, who prepared it in 1886, 
 dropped some of the liquid fluorine on the wooden floor of the laboratory 
 when the wood immediately burst into flame. Because of its remarkable 
 chemical activity, it is exceedingly difficult to prepare. Moissan pre- 
 pared it by electrolysis of pure liquid H 2 F 2 ( 25 C. to 50 C.), mixed 
 with a little KF. 
 
 H +F HF + 37,600 cal. 
 
THE NON-METALS, OR ACID-FORMING ELEMENTS. 153 
 
 Hydrogen Derivatives. 
 
 Experiment I. Preparation and Properties of Hydrogen Fluorine. 
 
 To a gram of powdered CaF 2 in a test tube add 3 cm. 3 of concentrated 
 H 2 SO 4 ; heat the tube gently; test the effect of the evolved gas upon a 
 piece of moistened blue litmus paper. Dip a glass rod into a few cm. 3 
 of NH 4 OH in a test tube; hold rod in evolved gas. Results? After 
 the test tube has been thoroughly cleaned, observe the effect of the gas 
 upon the inside surface of the glass. Write equations indicating nature 
 of above reactions. 
 
 Caution. The above experiment should be performed in the hood. 
 The gas must not come in contact with the skin nor be breathed. 
 
 Experiment II. Etching of Glass with Hydrofluoric Acid. 
 
 Spread a smooth thin layer of paraffin wax upon one side of a piece of 
 glass, about 10 cm. square; allow the wax to cool and harden; by means 
 of the sharp end of a file or any pointed instrument, cut through the wax 
 to the glass, making any chosen design. Cover the bottom of a shallow 
 lead dish with CaF 2 ; add sufficient warm H 2 SO 4 to moisten powder in 
 dish; place glass plate, waxed side downward, upon the lead dish; warm 
 dish gently but do not melt paraffin. After 10 min. remove heat, and 
 cool dish; remove glass plate and wash off the wax. Has the gas, H 2 F 2 , 
 evolved affected the glass plate ? Equation ? 
 
 Caution. Use the greatest care to avoid breathing the gas and to 
 prevent the acid from coming in contact with skin. Perform experiment 
 in the hood. 
 
 Experiment III. Formation of Silicon Tetrafluoride by the Action of 
 H 2 F 2 on Silica. 
 
 Mix i gram of silica, SiO 2 (sand), with i gram of CaF 2 ; place the 
 mixture in a test tube; add a little concentrated H 2 SO 4 ; heat gently. 
 Expose a drop or film of H 2 O to the action of the fumes of Si F 4 ; observe 
 the milky appearance of the water due to the formation of silicic acid, 
 H 4 Si0 4 . 
 
 (i) CaF 2 + H 2 S0 4 - 
 
 (2) SiO 2 + 2H 2 F 2 SiF 4 + 2H 2 O, 
 
 (3) 3 SiF 4 + 4 H 2 -> 2 H 2 Si F 6 + H 4 Si O 4 . 
 
 The fluorides of lithium, sodium, potassium, ammonium, silver, 
 mercury, iron, aluminum and tin are soluble in water. 
 
 CHLORINE Cl. 
 
 At. Wt. 35.45 Mol. Wt. 70.9. 
 Preparation and Properties oj Chlorine. 
 
 Experiment I. Review the previous experiments with chlorine. Fix 
 in mind the physical and chemical properties of the elementary substance. 
 Hydrogen Derivatives. 
 
154 EXPERIMENTAL CHEMISTRY. 
 
 Experiment II. Preparation and Properties of Hydrogen Chloride. 
 
 Recall the facts about hydrogen chloride. Compare its properties 
 with the corresponding hydrogen compounds of the other halogens. 
 These hydrogen compounds are styled the hydrogen halides. 
 
 H + Cl HC1 + 22,000 cal. 
 
 All chlorides, with the exception of silver, mercury (ous), lead and a 
 few basic chlorides like BiOCl and SbOCl are soluble in water. PbCl 2 
 is slightly soluble in cold water and easily soluble in hot water. 
 
 Oxygen Derivatives oj Chlorine. 
 
 Experiment III. Preparation of Sodium Hypochlorite and Hypo- 
 chlorous Acid. Oxidizing Power of Hypochlorous Acid Bleaching. 
 
 (a) Dissolve 3 grams of NaOH in 20 cm 3 , of H 2 O; pass chlorine into 
 the cold dilute solution, but do not saturate it. The probable reaction 
 may be indicated by the following equation: 
 
 2 NaOH + C1 2 - Nad + NaOCl + H 2 O. 
 
 (b) Divide the solution into two parts; to one part add dilute H 2 SO 4 ; 
 observe the odor; suspend a piece of cheap red calico and a strip of 
 litmus paper in the acidulated solution for a day; then observe the 
 bleaching effect of hypochlorous acid, HC1O upon paper and calico. 
 Write equations. 
 
 Ascertain the bleaching properties of a non-acidulated solution of 
 NaOCl by repeating above, using the second portion of the solution. 
 After removing paper and calico, add a few drops of the solution to an 
 indigo, C 16 H 10 N 2 O 2 , solution. Effects? Add a few cm. 3 of dilute 
 H 2 SO 4 to solution, then add 3 cm. 3 of the solution to the indigo; observe 
 the bleaching effect. Equations explaining activity of HC1 are as 
 follows: 
 
 2 HC1O 2HC1 + 2O + 18,600 cal. 
 C 16 H 10 N 2 O 2 + 2(3 2C 8 H 5 NO 2 + 1,800 cal. 
 
 C 16 H 10 N 2 O 2 + 2HC1O -> 2 C 8 H 5 NO 2 + 2 HC1 + 20,400 cal. 
 
 Experiment IV. Bleaching Powder ("Chloride of Lime"). 
 
 Place 5 grams of chloride of line (Ca(OCl) 2 + CaCl 2 ) in a flask; add 
 20 cm. 3 of H 2 O and shake thoroughly; filter. Using the clear filtrate, 
 repeat Exp. III. Do not neglect to acidulate solution or moisten objects 
 to be bleached with very dilute H 2 SO 4 (i drop in 25 cm. 3 of H 2 O). 
 
 Experiment V. Preparation and Properties of Potassium Chlorate. 
 
 (a) Dissolve 3 grams of KOH in 10 cm. 3 of H 2 O; saturate the boiling 
 solution with chlorine; cool the solution and observe the formation of 
 crystals of potassium chlorate, KC1O 3 . Devise a method for proving 
 
THE NON-METALS, OR ACID-FORMING ELEMENTS. 155 
 
 that this substance is identical with the KC1O 3 (solid) found upon the 
 end shelves. Equations ? 
 
 (b) KC1O 3 as an oxidizer. Mix a small quantity of powdered KC1O 3 
 with a little powdered charcoal; place the mixture upon the cover of a 
 crucible and heat gently. Results ? When the cover has cooled, place 
 it in a small beaker which contains 25 cm. 3 of H 2 O; allow the products of 
 the fusion to dissolve; filter; add a few drops of the filtrate to 2 cm. 3 
 of AgNO 3 ; if a white precipitate forms which turns dark, on exposure 
 to light it is AgCl. Explain by use of equations. 
 
 Experiment VI. Chloric Acid. Chlorine Dioxide. 
 
 (a) Place a small crystal of KC1O 3 in the bottom of a test tube; add 
 a few drops of strong H 2 SO 4 ; point the mouth of the tube away from 
 your face, then heat cautiously; the explosive action is due to the presence 
 of a yellow gas, chlorine dioxide, C1O 2 , which is violently explosive owing 
 to its ready decomposition into chlorine and oxygen, with liberation 
 of much heat. 
 
 2 KC10 3 + H 2 S0 4 2 HC10 3 + K 2 S0 4 , 
 
 4HC1O 3 - 2H 2 O + 4 C1O 2 + 2O, 
 2 HC1O 3 + 2 O -> 2 HC1O 4 ; 
 or 
 
 2C1O 2 CT 2 + 2O 2 ; 
 
 3HC1O 3 HC1O 4 + 2OO + H 2 O, 
 again, 
 
 6KC10 3 + 3 H 2 S0 4 -> 2 HC10 4 + 4 C1O 2 + 3K 2 SO 4 + 2 H 2 O. 
 Chloric acid Structural formula: H O Cl /^ Q 
 
 (b) Oxidizing action of chloric acid (C1O 2 ). 
 
 Powder i gram of KC1O 3 ; place the powder upon a sheet of paper 
 and mix with it i gram, of sugar; put the mixture upon an iron plate; 
 allow a drop of strong H 2 SO 4 to drop from the end of a glass rod upon 
 the mixture. Result? 
 
 Note. White gunpowder is a mixture of KC1O 3 , K 4 Fe(CN) 6 and 
 C i2H 22 O n . 
 
 HC1O 3 , Aq > HC1, Aq + 30 + 15,300 cal. 
 
 (c) Explosive conduct of chlorates. Place very small quantities of 
 KC1O 3 and sulphur in a mortar; rub the mixture vigorously with the 
 pestle. Sharp explosions result from the friction. As a rule, should 
 
 . chlorates be pulverized with other substances ? 
 
 All chlorates are soluble in water. 
 
 Experiment VII. Perchlorates and Perchloric Acid. 
 
 Perchlorates may be formed by suitably heating chlorates; corre- 
 sponding chlorides will be formed simultaneously. When either NaClO 3 
 or KC1O 3 is heated oxygen is evolved; if the operation is stopped when 
 
156 EXPERIMENTAL CHEMISTRY. 
 
 one-third of the oxygen has been liberated, the chloride and the chlorate 
 of the metal remain in the fused mass. The more soluble chloride may 
 be dissolved out from the mass. 
 
 2 KC10 3 KC10 4 _+ KC1 + Q; 
 KC10 4 -*KC1 + 20 2 . 
 
 Perchloric acid, HC1O 4 , may be formed by the action of strong H 2 SO 4 
 on perchlorates. Equation? 
 
 HC1O 4 , Aq HC1, Aq + lzO~ 2 + 700 cal. 
 Structural formula for HC1O 4 ? 
 
 BROMINE, Br. 
 
 At. Wt. 79.96 Mol. Wt. 159.92. 
 Preparation and Properties oj Bromine. 
 Experiment I. (Hood.) 
 
 Mix 2 grams of powdered MnO 2 with an equal weight of pulverized 
 NaBr or KBr. Introduce the mixture into the test tube used when pre- 
 paring oxygen from KC1O 3 and MnO 2 ; add a few cm. 3 of strong H 2 SO 4 ; 
 heat gently and conduct the gas into a 50 cm. 3 flask which is nearly im- 
 mersed in cold water. Observe color, odor and relative density of the gas. 
 When one or two drops of the bromine gas have condensed in flask, 
 conduct the bromine into a test tube containing 15 cm. 3 or 20 cm. 3 of 
 H 2 O. Compare the drop of condensed bromine vapor with the liquid 
 bromine found on side shelf. If the H 2 O in the test tube was not satu- 
 rated with bromine, add a drop of the liquid bromine to the bromine 
 water (H 2 -f Br.); shake thoroughly. What is the color of the solution 
 (bromine water) ? Test the action of the solution on litmus paper. To 
 one portion of the solution add 2 cm. 3 of ether; to a second portion 2 cm. 3 
 of carbon disulphide; and to a third portion, 2 cm. 3 of chloroform; shake 
 each tube vigorously. Results? 
 
 The following partial equations may assist in interpreting the reactions 
 involved in preparing bromine: 
 
 NaBr + H 2 SO 4 -> Na 2 HSO 4 + (H Br). 
 MnO 2 + H 2 SO 4 MnSO 4 -KH 2 O + (O). 
 2HBr + O < H 2 O + Br 2 . 
 or 
 
 4HBr + MnO 2 MnBr 2 + 2 H 2 O + BF 2 . 
 MnBr 2 + H 2 SO 4 MnSO 4 + 2HBr. 
 
 Note. If larger quantities of bromine are to be prepared a retort and 
 a receiving flask may be substituted for the above apparatus. 
 
 Experiment II. Substituting Power of Chlorine as Compared with 
 Bromine. 
 
 (a) Dissolve a very small crystal of NaBr or KBr in 15 cm. 3 of water in 
 
THE NON-METALS, OR ACID-FORMING ELEMENTS. 157 
 
 a test tube; add a few cm. 3 of chlorine water (side shelf); divide the 
 solution into two parts. Heat one portion gently and note the change 
 in color. Explain. 
 
 (b) To the second portion add 2 cm. 3 of CS 2 ; shake thoroughly. 
 Results? Equations? Which of the two chlorine or bromine has 
 the greater affinity for sodium ? 
 
 Exp. III. Displacement of Sulphur from Hydrogen Sulphide. 
 
 Prepare a saturated solution of bromine by adding a drop of liquid 
 bromine to 10 cm. of H 2 O; pass H 2 S through the solution until the brown 
 color due to the presence of bromine has disappeared. Is there any 
 evidence of the presence of free sulphur? Filter the solution; boil the 
 nitrate until the fumes will not darken a piece of filter paper moistened 
 with a solution of lead nitrate; test the nitrate with the litmus paper. 
 To 2 cm. 3 of AgNO 3 add H 2 S; repeat, using a few drops of the above 
 nitrate. Equations for all reactions? What substance is produced 
 by the displacement of sulphur from H 2 S by bromine? Recall the 
 similar conduct of chlorine in an analogous experiment. 
 
 Hydrogen Derivatives. 
 
 Experiment IV. (Hood.) Preparation and Properties of Hydrogen 
 Bromide Preliminary. 
 
 Pulverize about 3 grams of NaBr or KBr, and place it in a test tube; 
 add sufficient strong H 2 SO 4 to thoroughly moisten powder; heat gently; 
 blow your breath across the mouth of the tube. Results ? Dip a glass 
 rod into NH 4 OH, then hold rod near mouth of test tube. Results? 
 Hold a moistened piece of blue litmus paper in the evolved gas. Results ? 
 What is the color of hydrobromic acid, HBr? 
 
 What hydracids studied previously conduct themselves in a similar 
 manner toward the identical tests? Write the equations for the above 
 reactions. 
 
 NaBr + H 2 SO 4 NaHSO 4 + HBr. 
 
 What is the effect of increasing the temperature of the mixture in the 
 test tube ? What is the colored gas which makes its appearance ? 
 
 NaBr + H 2 SO 4 NaHSO 4 _+ (HBr), 
 2HBr + H 2 SO 4 -> SO 2 + Br 2 + H 2 O. 
 
 H 2 SO 4 acts as an oxidizing agent. Was either free fluorine or chlorine 
 observed during the preparation of their respective hydracids from 
 corresponding salts by the use of H 2 SO 4 ? 
 
 Experiment V. Laboratory Source of Hydrogen Bromide. Hydro- 
 bromic Acid. 
 
 Into a test tube which has been provided with a stopper and a de- 
 livery tube, introduce 5 grams of powdered NaBr or KBr; add sufficient 
 concentrated phosphoric acid, H 3 PO 4 , to cover the powder; shake until 
 
158 
 
 EXPERIMENTAL CHEMISTRY. 
 
 the two substances are mixed thoroughly; fit stopper and delivery tube 
 into place; heat tube gently; collect the evolved gas in 5 cm. 3 of H 2 O 
 in a test tube the mouth of the delivery tube should be about .2 cm. 
 above the surface of the water. 
 
 Is there any evidence of the presence of free bromine? In what 
 particular respect is H 2 SO 4 different from H 3 PO 4 in its action on bro- 
 mides ? The acid solution may be tested as in Experiment IV. Add 
 a few drops of the solution to a little sodium carbonate. Results? 
 Repeat the foregoing experiment, using a few cm. 3 of AgNO 3 . Results ? 
 
 FIG. 28. 
 
 If the acid solution is saturated, its effect on a piece of zinc should be 
 ascertained. Wrap a piece of platinum w r ire around the zinc. Why? 
 Equations ? 
 
 H + Br -* HBr + 8400 cal. 
 Experiment VI. Optional. Preparation of Pure Hydrobromic Acid. 
 
 Note. When phosphorous and bromine are mixed, they combine 
 energetically forming phosphorous tribromide (PBr 3 ). Pure HBr is 
 prepared by hydrolysis of this substance. 
 
 /Br H OH 
 
 /OH 
 
 P Br H OH P OH + 3 H Br. 
 \Br-H OH \OH 
 
 The preparation of the acid may be conducted as follows : The ordinary 
 funnel of a 300 cm. 3 generating flask is displaced by a dropping funnel; the 
 flask is connected in series with two U- tubes (or see Fig. 28) ; the U-tube next 
 to the flask should be half filled with glass beads which have been moist- 
 ened and rolled in red phosphorous; the second tube should be about 
 one-fourth filled with water to absorb the HBr. The mouth of the tube 
 
THE NON-METALS, OR ACID-FORMING ELEMENTS. 159 
 
 delivering the gas to the tube should be above the surface of the water. 
 Charge the generating flask with 5 grams of red phosphorus and enough 
 water to barely cover the latter. (It is suggested that 5 grams to 10 grams 
 of clean sand be added with the charge.) Pour 10 cm. 3 of bromine 
 (liquid) into the funnel; when apparatus has been assembled properly, 
 allow the bromine to fall on the phosphorus drop by drop. The union 
 of the bromine and phosphorus will be accompanied by a flash of light. 
 (Can you suggest a reason for using red instead of yellow phosphorus ?) 
 After the water in the flask has been saturated with the HBr, the latter 
 will escape into the U-tube and eventually be absorbed in the water con- 
 tained in the most remote U-tube. If too little water is present in the flask, 
 crystals of PBr 3 will clog the apparatus; the addition of a few drops of 
 water will prevent their formation. The acid properties of the solution 
 of the gas may be tested as previously. If convenient its conductivity 
 may be -determined. 
 
 The bromides are easily soluble' in water with the exception of silver, 
 mercury (ous) and lead. PbBr 2 is slightly soluble in cold water and 
 easily soluble in hot water. 
 
 Oxygen Derivatives oj Bromine. 
 
 Note. No oxides of bromine have been prepared. 
 
 Experiment VII. Preparation of Sodium Hypobromite and Hypo- 
 bromous Acid. 
 
 (a) Dissolve 3 grams of NaOH in 20 cm. 3 of H 2 O; add a few drops of 
 liquid bromine to the cold dilute solution. Observe the disappearance 
 of the color of the bromine. The bromine should be added drop by 
 drop, but not in a sufficient quantity to saturate the solution. 
 
 2 NaOH + Br 2 NaBr + NaBrO + H 2 O. 
 
 (b) Divide the above solution into two parts. To one portion add 
 dilute H 2 SO 4 ; observe the color and odor of the evolved gas. The 
 following partial equations probably represent the stages of the reaction: 
 
 NaBr + NaBrO + H 2 SO 4 -* Na 2 SO 4 + (HBr) + (HBrO). 
 HBr + HBrO H 2 O + Br 2 . 
 Adding, NaBr + NaBrO + H 2 SO 4 -> Na^SO, + H 2 O + Br 2 . 
 
 (c) Repeat (b) using acetic acid instead of H 2 SO 4 . Use slight excess 
 of acid. Note that bromine is liberated. Write the structural formula? 
 for NaBrO and NaBrO 3 . 
 
 Experiment VIII. Preparation and Properties of Potassium Bromate. 
 
 Dissolve 3 grams of KOH in 10 cm. 3 of H 2 O; saturate the solution 
 with bromine; heat to boiling. When crystals make their appearance, 
 dissolve them in H 2 O; add a few drops of liquid bromine; boil to crystal- 
 lization; dissolve crystals in the smallest quantity of water; pour the 
 
l6o EXPERIMENTAL CHEMISTRY. 
 
 solution into a crystallizing dish. The KBrO 3 crystals may be separated 
 from the more soluble KBr crystals by crystallization. Equation? 
 
 Experiment IX. Bromic Acid. Interaction of Bromic and Hydro- 
 bromic Acids. 
 
 (a) To a solution of NaBrO 3 or KBrO 3 , add dilute H 2 SO 4 . Is bro- 
 mine liberated? Equation? 
 
 (b) To a solution containing KBrO 3 add dilute H 2 SO 4 . Is bromine 
 liberated ? Equations ? 
 
 The following thermo-chemical equation gives an approximate idea 
 of the oxidizing power of bromic acid: 
 
 HBrO 3 , Aq * HBr, Aq + 30 + 15,000 cal. 
 Structural formula for hydrobromic acid: 
 
 O 
 
 S 
 
 H O Br 
 
 \ 
 O 
 
 IODINE. I. 
 
 At. Wt. 126.97 Mol. Wt. 254.8 (i85-6oo). 
 Preparation and Properties oj Iodine. 
 
 Experiment I. 
 
 (a) Mix 2 grams of powdered MnO 2 with an equal weight of pulverized 
 Nal or KI; place the mixture in an evaporating dish; add sufficient 
 concentrated H 2 SO 4 to moisten the mixture thoroughly; clamp an in- 
 verted funnel over the evaporating dish in such a position that all vapors 
 from the dish will escape through the stem of the funnel. Heat the 
 dish gently. A dense vapor will be set free. Observe color, odor and 
 relative density of the vapor before it condenses upon the sides of the 
 funnel. Continue to heat dish until colored vapor ceases to be evolved. 
 After cooling funnel, scrape crystals into a clean beaker to be used for 
 following experiments. 
 
 (b) Resublimed iodine. Place a few of the crystals in a flask; heat 
 the bottom of the flask gently to volatilize the iodine which crystallizes 
 upon the cooler portions of the flask. Does the iodine fuse before be- 
 coming a vapor? What are such processes called? 
 
 (c) Test the solubility of iodine in each of the following solvents: 
 H 2 O, C 2 H 5 OH, CHC1 3 , (C 2 H 5 ) 2 O, CS 2 , and a KI solution. Recall a 
 similar experiment under subject of " Solutions." 
 
 (d} Triturate .3 gram of starch and 20 cm. 3 of H 2 O in a mortar until 
 they are thoroughly mixed. Pour this mixture into 100 cm. 3 of boiling 
 water; boil the mixture for about five minutes. Add a few drops of the 
 cold starch emulsion to a few cm. 3 of a KI solution of iodine; heat the 
 mixture. Does the color disappear ? Cool. Results ? The compound 
 
THE NON-METALS, OR ACID-FORMING ELEMENTS. l6l 
 
 formed with the characteristic blue color is known as starch iodide. 
 Iodine is used to detect the presence of starch^ and vice versa. 
 
 Note. Bottle the starch emulsion and preserve it for future use. 
 
 The solution of iodine in alcohol is known as the tincture of iodine. 
 The brown color of the tincture is attributed to the fact that the iodine 
 and alcohol are in a loose state of combination; likewise iodine dis- 
 solved in a KI solution is supposed to form a definite compound: 
 
 2 KI + I 2 ~* 2KI 3 . 
 
 On the other hand, the blue substance, starch iodide, referred to as a 
 compound, is not a^ chemical compound but merely a solution of iodine 
 in starch. 
 
 Experiment II. Substituting Power of Chlorine and Bromine as 
 Compared with Iodine. 
 
 (a) Dissolve a very small crystal of Nal or KI in 15 cm. 3 of water in 
 a test tube; divide it into two portions; to one portion add a few drops of 
 chlorine water; heat gently and observe the change in color; add a few 
 drops of the solution to a little of the starch solution. Results ? Shake 
 up the rest of the solution with 3 cm. 3 of CS 2 . Results? Equations? 
 
 Note. An excess of chlorine water decolorizes a solution of iodine 
 in CS 2 . 
 
 (b) Repeat (a) using the second portion of the solution and bromine 
 water instead of chlorine water. Results ? Equations ? 
 
 In view of the above and previous experiments what do you infer in 
 regard to the relative affinities of the halogens for hydrogen? Arrange 
 them in the order of their increasing affinity. 
 
 Experiment III. Potassium Bichromate Liberates Iodine from an 
 Acidified Solution of an Iodide. 
 
 To a dilute solution of KI add CS 2 ; shake. Results? Acidify the 
 solution with a little dilute H 2 SO 4 , and add a few drops of a K 2 Cr 2 O 7 * 
 solution; agitate the mixture. Results? 
 
 Note, Bromine is not liberated from bromides by above process. 
 Test. Equations ? 
 
 Experiment IV. Iodine is Liberated by the Oxidizing Action of Sul- 
 phuric Acid. 
 
 Place a small crystal of Nal or KI in a test tube; add 3 cm. 3 of concen- 
 trated H 2 SO 4 ; heat gently. Observe the color and odor of the evolved 
 gas. Identify it. Ascertain whether hydriodic acid, HI, is evolved or 
 not. 
 
 KI + H 2 S0 4 <= KHS0 4 + (HI), 
 2 HI + H 2 S0 4 ^ S0 2 + I 2 + H 2 0, 
 6HI + H 2 S0 4 <= S + 3 I 2 + 4 H 2 0, 
 SHI + H 2 SO 4 * H 2 S + 4l 2 + 4H 2 O. 
 * Sodium nitrite may be sustituted for K 2 Cr 3 O 7 . 
 ii 
 
1 62 EXPERIMENTAL CHEMISTRY. 
 
 Compare the above reaction with the action of H 2 SO 4 on CaF 2 , NaCl, 
 NaBr. Arrange the hyrdascids of the halogens in the order of their 
 increasing stability toward oxidizing agents. 
 
 Experiment V. Displacement of Sulphur from Hydrogen Sulphide. 
 
 To 2 grams of powdered iodine in a flask fitted with a cork containing 
 a single perforation add 20 cm. 3 of H 2 O; pass H 2 S through the mixture 
 until iodine disappears. Proceed with experiment as per directions 
 Exp. Ill, "Bromine." 
 
 Hydrogen Derivatives. 
 
 Experiment VI. Preparation and Properties of Hydrogen Iodide- 
 Hydriodic Acid. 
 
 (a) Follow directions given in Exp. V., "Bromine," substituting the 
 words iodide and iodine for bromide and bromine, respectively. Equa- 
 tions ? 
 
 (b) Pure HI may be prepared by a process similar to that used in the 
 preparation of pure HBr Exp. VI., "Bromine." Use the same appa- 
 ratus. The flask may be charged with a mixture of iodine and water 
 (4 to i), and the red phosphorus (stirred to a paste with water) allowed 
 to fall slowly on the iodine, drop by drop. 
 
 Another method, using same apparatus, is as follows: Mix the red 
 phosphorus and iodine (10 to i by weight) in a dry mortar; charge flask 
 with this mixture, and allow water to fall slowly upon it drop by drop. 
 H -f I HI 6000 cals. 
 
 Experiment VII. (Optional.) Use of Potassium Iodide to Determine 
 Rate of Absorption from the Stomach. 
 
 Test the saliva for iodine by the iodide of starch test. Take, by mouth, 
 a capsule containing 0.2 gram of KI or 10 drops of the saturated solution. 
 Note the time. Test the saliva at end of each minute for iodine. If 
 absorption is normal, a positive test should appear in fifteen minutes. 
 Taking equal quantities of the iodide, the members of the class will be 
 able to make a comparison of the rates of absorption. Determine the 
 average rate of absorption of the individuals of the class who are fleshy; 
 likewise calculate the average rate of those who are not fleshy. Your 
 conclusions ? 
 
 DERIVATIVES OF IODINE. 
 
 Oxygen. 
 
 Note. Although iodine and ordinary oxygen do not combine with 
 one another, O 3 reacts with iodine to form I 2 O 5 . Compare with con- 
 duct of fluorine. 
 
 Experiment VIII. Preparation of Sodium Hypoiodite. Sodium lodate. 
 
 These salts may be prepared by a method analogous to that used in 
 the preparation of the corresponding salts of bromine and chlorine. 
 
THE NON-METALS, OR ACID-FORMING ELEMENTS. 163 
 
 The hypoiodite, however, oxidizes quite readily, forming the iodate and 
 iodide. If this mixture is treated with acid, iodic and hydriodic acids 
 are liberated, but interact, liberating iodine. Recall the conduct of 
 corresponding acids of bromine. The oxacids of iodine, HIO and HIO 3 , 
 are quite stable. Hydriodic acid: 
 
 O 
 
 ^ 
 H O I 
 
 \ 
 O. 
 
 Experiment IX. Relative Affinity of Bromine and Iodine for Oxygen. 
 
 To 5 cm. 3 of a solution of KBrO 3 add 5 cm. 3 of dilute H 2 SO 4 . 
 Divide the solution into two parts. To one portion add CS 2 and shake. 
 Is the CS 2 colored? To the other portion add a very small flake of 
 iodine; shake vigorously for two or three minutes; decant the clear 
 solution into another test tube; add a few drops of CS 2 ; agitate the con- 
 tents. Does the reddish-brown color indicate that iodine has a greater 
 affinity than bromine for oxygen ? Your reasons ? (Examine structural 
 formula of HBrO 3 .) 
 
 Experiment X. Properties of lodates. 
 
 (a) Heat a few crystals of KIO 3 or NaIO 3 in a test tube and test for 
 oxygen. Recall the preparation of oxygen from KC1O 3 . Equation? 
 
 (b) Pulverize .5 gram of KIO 3 or NaIO 3 ; mix thoroughly with .3 gram 
 of powdered charcoal; place the mixture in a crucible and heat until the 
 entire mass glows; cool; place crucible and contents in a beaker of hot 
 water to dissolve the fused mass. Filter the solution. Devise a method 
 to prove that an iodide was formed by the above reaction. 
 
 2NaIO 3 + 3C 2NaI + 3CXX. 
 
 (c) Place a small crystal of an iodate in a mortar with a little sulphur; 
 rub the mixture vigorously with a pestle. Sharp explosions result from 
 the friction. What other substance reacts like the iodate ? Equations ? 
 Would you infer that iodates are strong oxidizers? Structural formula 
 of NaIO 3 ? 
 
 Experiment XI. Preparation of Iodic Acid. 
 
 Note. Iodic acid may be prepared by a method analogous to the 
 preparation of chloric acid, i.e., by the action of H 2 SO 4 on a soluble 
 iodate. 
 
 2 NaIO 3 + H 2 SO 4 Na 2 SO 4 + 2 HIO 3 . 
 
 The iodic acid may be crystallized out from the solution as a white 
 solid. 
 
 The author has found the following method for the preparation of 
 the acid to be very satisfactory: Place two or three small flakes of 
 
164 EXPERIMENTAL CHEMISTRY. 
 
 iodine in the bottom of the largest test tube'among your apparata; add 
 15 cm. 3 of strong HNO 3 ; take the mixture to the hood and boil vigorously 
 until the solution becomes colorless. It may be necessary to add HNO 3 
 from time to time to oxidize all of the iodine. The colorless solution of 
 HIO 3 should be boiled for sometime to concentrate it, and then set aside 
 to allow crystals of the acid to form, or the solution may be tested at 
 once as follows: 
 
 Add a few drops of HNO 3 to an AgNO 3 solution. Does a precipitate 
 form ? Repeat foregoing using a solution of KI instead of the HNO 3 ; ob- 
 serve the yellow color of the precipitate, Agl. Now add a few drops of 
 the prepared solution supposed to contain iodic acid, to 3 cm. 3 of the 
 AgNO 3 solution. What is the color of the precipitate ? Was it formed 
 by the action of HNO 3 ? Your reasons ? 
 
 3 -* 2H 2 O + 4NO 2 + (O 2 ). 
 I 2 + H 2 O + 50 -> 2 HIO 3 . 
 
 Save the iodic acid solution for the following experiment. 
 Experiment XII. Oxidizing Action of Iodic Acid. 
 
 To 5 cm. 3 of the solution of iodic acid add 2 cm. 3 of CS 2 ; mix thoroughly. 
 Note that the color of the CS 2 remains unchanged. Add a few drops 
 of sulphurous acid, H 2 SO 3 to the solution and then shake thoroughly. 
 What is the substance which imparts the color to the CS 2 ? Has the 
 odor of SO 2 disappeared? 
 
 2 HIO 3 + 5H 2 SO 3 -> sH 2 SO 4 + I 2 + H 2 O. 
 HIO 3 , Aq HI, Aq + 36 42,900 cal. 
 
 Experiment XIII. (L. T.) Periods of Induction in Chemical Reac- 
 tions. 
 
 Instructions. "There is a pleasing 'lecture experiment' for illus- 
 trating the 'period of induction.' A very dilute solution of H 2 SO 3 
 and HIO 3 (i gram, e.g., in 600 liters of water) is mixed with starch. The 
 appearance of a visible blue color occupies a measurable time, which 
 may be extended by using more dilute solutions." Mellor's "Chemical 
 Statics and Dynamics." 
 
 Experiment XIV. Resemblance of the Properties of Corresponding 
 Compounds of the Halogens. 
 
 (a) To a solution of Pb(NO 3 ) 2 add a few drops of a KC1 solution; 
 observe color of precipitate; filter. Determine solubility of precipitate 
 in hot water. Repeat above using KBr. Also KI. 
 
 (b) Repeat (a) substituting AgNO 3 for Pb(NO 3 ) 2 , and NH 4 OH for 
 hot water. 
 
 (c) Repeat (b) substituting HgNO 3 for AgNO 3 . 
 
 Note the general similarity of conduct of the respective halides of each 
 metal. Tabulate results. 
 
THE NON-METALS, OR ACID-FORMING ELEMENTS. 
 
 Experiment XV. Detection of Fluorides, Chlorides, Bromides and 
 Iodides together in a Solution. 
 
 Devise a system of tests which will provide for the detection of the 
 halogens in solution. Apply the system to a solution containing the 
 four halides. Ask the assistant to give you an " unknown" solution; 
 test for the presence of the halides. Report ( ?). Make a complete record 
 of all work. 
 
 A general outline of the relations of the halogens is presented in the 
 following table: 
 
 Physical Properties Fluorine Chlorine Bromine Iodine 
 
 Atomic w r eight 19 35-45 79-9^ 126.85 
 
 State or phase Gas Gas Liquid Solid 
 Color 
 
 Specific gravity 
 Specific heat 
 Molecule of gas 
 Melting point 
 Boiling point 
 Chemical Properties 
 Water decomposed 
 
 Fluorine 
 
 Chlorine Brc 
 
 Gas 
 
 35-45 7^ 
 Gas Lii 
 
 Pale 
 
 Greenish- 
 
 yellow 
 15 (liquid) 
 
 yellow Brc 
 1.15 (liquid) 3. 
 
 Black 
 
 F 2 C1 2 
 
 102 (solid) 
 
 186 
 
 -70 >7 
 
 33-7 
 
 Br 2 I 2 (2oo-6oo) 
 
 o 
 l8 4 
 
 7-3 H4 
 
 59 
 
 In the dark, In the sun- Slowly in Only by in- 
 light, the sun- direct pro- 
 light, cesses. 
 
 Readily, In the sun- In heat of Only by in- 
 light, flame, direct pro- 
 cesses. 
 
 H F HC1 H Br HI 
 
 37,600 cal. 22,000 cal. 8400 cal. 6000 cal 
 
 Formation of H- 
 derivatives 
 
 Heat of formation 
 of H-derivatives 
 
 Stability of H- 
 derivatives 
 Affinity for oxygen 
 Oxides 
 
 Oxacids 
 
 V 
 
 I* 
 V 
 
 * The stability of the oxacids and their corresponding salts increases with increase 
 jn oxygen-content. 
 
 ^v 
 
 X, 
 
 
 
 v 
 
 s 
 
 None 
 None, 
 
 C1 2 0,C1 2 3 , 
 C1O 2 . 
 HC10 
 HC1O 2 
 HC1O 3 
 HC10 4 
 
 
 IA 
 
 HTCr O 
 
 
 
 H Br O 3 H 
 H Br O 4 ? H 
 
 10, 
 10. 
 
1 66 EXPERIMENTAL CHEMISTRY. 
 
 It must be obvious to the student that the elements of this group or 
 family are connected by properties, physical and chemical, which show 
 a remarkable similarity. As a matter of fact, the uniformity is so marked 
 that if the elements are arranged so as to show a gradation of any one 
 property, the order will not be altered if they are arranged with respect 
 to other properties. Again, these properties seem to bear so definite a 
 relation "to the atomic weight of an element, that it is now generally 
 believed by chemists, although it is not yet fully proved, that the proper- 
 ties of the elements, as well as the properties of their compounds, are 
 periodic junctions oj their atomic weights. For example, the specific heat 
 of an element is inversely proportional to its atomic weight. These and 
 other facts have led to the grouping of the elements into " natural families " 
 or groups. The halogens are unique in the fact only that they represent 
 the most prominent example of the gradation in properties observed 
 among the members of a natural group of elements. 
 
 The regular, periodic variation in the properties of the elements con- 
 stitutes what is known as the periodic law, by means of which the elements 
 are given their natural classification. The law is illustrated by the 
 following tables. (See Classification of the Elements.) 
 
 It should be noted before leaving this group that the halogen elements 
 are, par excellence, acid elements; not only do their oxygen compounds, 
 but their hydrogen compounds also exhibit acid properties. 
 
CHAPTER XVI. 
 CLASSIFICATION OF THE ELEMENTS. 
 
 The classification of the elements into metals, non-metals and metal- 
 loids was made ^during the early development of chemistry when com- 
 paratively few elements were known. This classification was convenient, 
 but it soon led to confusion, because elements were discovered which 
 did not fall exclusively into any one of these three groups. It soon 
 became evident to the chemists of that time that a new plan of classi- 
 fication was necessary; especially so after the announcement of the 
 "Law of Multiple Proportion" by Dalton. As the result of the accept- 
 ance of this law, Prout (1815) promulgated the theory that the atomic 
 weights of all the elements are exact multiples of the atomic weight of 
 hydrogen, assuming that the atoms of other elements are merely con- 
 densations or aggregations of hydrogen atoms. The atomic weight 
 determinations of Berzelius, Marignac and Stas did not prove that the 
 atomic weights of the elements are exact multiples of the atomic weight 
 of hydrogen. 
 
 Note. " The work of T. W. Richards on the atomic weights of a 
 large number of the metals should receive special attention. He has 
 improved old methods, devised new ones and applied them with a skill 
 which is rare. His determinations are to be ranked among the very 
 best which have ever been made." Jones' "Physical Chemistry." 
 
 The theory of Prout, though proven untenable, was instrumental in 
 directing a great deal of attention to the question of atomic weights, with 
 the result, that Dobereiner (1825) "on examining the atomic weights 
 of correlated elements, observed that the atomic weight of the middle 
 member of a group of three related elements was nearly a mean of the 
 atomic weights of the other two elements. These groups of three were 
 known as the "Triads of Dobereiner" and probably represent the first 
 attempt to classify the elements on the basis of a relationship existing 
 between properties and atomic weights. The student is referred to 
 Venable's "Development of the Periodic Law." 
 
 NON-METALS. 
 
 Helium Fluorine Oxygen Nitrogen Carbon 
 
 Neon Chlorine Sulphur Phosphorus Silicon 
 
 Argon Bromine Selenium Arsenic Germanium* 
 
 Krypton Iodine Tellurium Antimony* Tin* 
 
 Xenon Bismuth* Boron 
 
 (Aluminium) 
 * Metalloids. 
 
 167 
 
i68 
 
 EXPERIMENTAL CHEMISTRY. 
 
 METALS. 
 
 See table of elements for complete list of metals, non-metals and metal- 
 loids. 
 
 TRIADS OF DOBEREINER (1825). 
 
 (Examples). 
 
 Lithium 
 
 Sodium 
 
 7 
 
 23 
 
 Potassium 
 
 Rubidium 
 
 39 
 
 8S 
 
 Calcium 
 
 Strontium 
 
 40 
 
 88 
 
 Sulphur 
 32 
 
 Selenium 
 
 77 
 
 Potassium 
 39 
 
 Caesium 
 133 
 
 B arium 
 
 137 
 
 Tellurium 
 127 
 
 7 + 39 
 
 = 23 
 
 39 + 133 
 
 40 + 137 
 
 86 
 
 3 2 + 127 
 
 = 78 
 
 Chlorine 
 
 35 
 
 Bromine 
 
 80 
 
 Iodine 
 126 
 
 35 
 
 = 80 
 
 H. 
 Li. 
 
 1 F. 
 
 2 Na. 
 
 8 Cl. 
 
 9 K. 16 Cu. 23 Rb. 
 G. 3 Mg. 10 Ca. 17 Zn. 24 Sr 
 Bo. 4 Al. 
 C. 5 -Si. 
 N. 6 P. 
 O. 7 S. 
 
 (Newland's, 1864.) 
 
 Co. & 
 15 Ni. 22 Br. 29 Pd. 36 I. 
 
 30 Ag. 37 Cl. ^ _. J0 
 3iCd.38Ba. &V.45 Pb. 54 
 
 _ TT T . T'l- ^ 
 
 Pt. & 
 
 42 Ir. 50 
 44 Tl. 53 
 
 10 ^a. 17 z,n. 24 sr. 31 (_,a. 30 &&. &V-45 Pb. 
 
 11 Cr. 18 Y. 25 Ce. &La. 32 U. 40 Ta. 46 Th. 56 
 
 12 Ti. 19 In. 26 Zr. 33 Sn. 39 W. 47 Hg. 52 
 
 13 Mn. 20 As. 27 Di. &Mo. 34 Sb. 41 Nb. 48 Bi. 
 
 14 Fe. 21 Se. 28 Ro. & Ru. 35 Te. 43 Au. 
 
 49 Os. 
 
 55 
 5 1 
 
CLASSIFICATION OF THE ELEMENTS. 
 
 169 
 
 
 
 
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 EXPERIMENTAL CHEMISTRY. 
 
CLASSIFICATION OF THE ELEMENTS. 
 
 171 
 
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172 
 
 EXPERIMENTAL CHEMISTRY. 
 
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CLASSIFICATION OF THE ELEMENTS. 
 
 173 
 
 Many attempts have been made to so modify the form of the fore- 
 going tables that the position of an element will more faithfully represent 
 its relationships. The following is an example: 
 
 BAYLEY S TABLE. 
 
 The following tabulated data may serve to make more significant 
 the relative positions of the elements in the periodic table: 
 
 Specific Gravity. 
 Na Mg Al Si P S Cl 
 
 .97 I.y 2.5 2.4 2.2 2.0 1.3 
 
 K. Ca Sc Ti V. Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br 
 
 .87 1.6 3.8 4. ? 5.5 6.8 7.2 7.9 8.5 8.9 8.9 7.2 5.9 5.5 5.6 4.8 3.2. 
 
174 
 
 EXPERIMENTAL CHEMISTRY. 
 
 
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CLASSIFICATION OF THE ELEMENTS. 
 
 ATOMIC VOLUME SPECIFIC GRAVITY. 
 
 Element. 
 
 ii 
 
 <u 
 
 < 
 
 Specific 
 Gravity.* 
 
 Atomic 
 Volume.* 
 
 Element. 
 
 'g bo 
 o'u 
 
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 -59 
 
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 8.9 
 
 7-i3 
 
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 Gas 
 
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 0.87 
 
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 16.5 
 
 
 
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 1.52 
 
 56-3 
 
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 0.97 
 
 2 3-7 
 
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 6.8 
 
 7-65 
 
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 2-5 
 
 34-4 
 
 Mg 
 
 24.0 
 
 1.74 
 
 13.8 
 
 Mn 
 
 5-50 
 
 7.14 
 
 7.6 
 
 Ag 
 
 107.9 
 
 10.6 
 
 10.2 
 
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 27.0 
 
 2.60 
 
 10.6 
 
 Fe 
 
 55-9 
 
 7.86 
 
 6-9 
 
 I 
 
 126.9 
 
 4-95 
 
 2 5 .6 
 
 Si 
 
 28.0 
 
 2 -39 
 
 10.7 
 
 Ni 
 
 58.7 
 
 8.90 
 
 6-31 
 
 Te 
 
 127 .6 
 
 6.25 
 
 20-4 
 
 P 
 
 31.0 
 
 2 . 20 
 
 12.8 
 
 Co 
 
 59-o 
 
 8-5 
 
 6.82 
 
 Cs 
 
 132.9 
 
 1.9 
 
 70. 
 
 *Huth's, Das periodische Gesetz der Atomgewichte. 
 
 EXERCISES. 
 
 1. Draw a periodic curve of the elements named in the above table, 
 tabulating the specific gravities on the vertical axis, and the atomic 
 weights on the horizontal axis. 
 
 2. Repeat (i), substituting atomic volumes for specific gravities. 
 
 3. Repeat (i), substituting specific heats (see Appendix) for specific 
 gravities. 
 
CHAPTER XVII. 
 OXYGEN FAMILY. 
 
 Oxygen, O, 16.00, 
 
 Sulphur, S, 32.06, 
 
 Selenium, Se, 79.2, 
 
 Tellurium, Te, 127.6. 
 
 It has been pointed out in the preceding chapter that the chemical 
 and physical properties of the elements bear a definite relation to their 
 atomic weights. The student should note carefully the conduct of these 
 four elements whose atomic weights place them in the same "natural 
 group," to ascertain whether they conform to this so-called "periodic 
 law." 
 
 OXYGEN, O. 
 
 At. Wt. 16. Mol. Wt. 32. 
 Preparation and Properties of Oxygen. 
 
 Experiment I. Recall the previous experiments with oxygen. Fix in 
 mind the physical and chemical properties of the elementary substance. 
 
 Hydrogen Derivatives. 
 
 Experiment II. Recall experiments suggested by the following equa- 
 tions: 
 
 H 2 + O > HO, Aq + 68,360 cals. 
 H 2 + 2 * H 2 2 > A q + 45.300 cals. 
 
 Experiment III. Oxidizing Action of Hydrogen Dioxide. 
 
 Place 5 cm. 3 of lead acetate in an evaporating dish; moisten a strip 
 of writing paper with the solution, then hold the paper in a current of 
 hydrogen sulphide. Results? Allow the paper to dry. Place 2 cm. 3 
 of H 2 O 2 in a test tube; using a glass rod wet with the H^Oj, write your 
 name upon the paper. Explain. 
 
 Oxygen Derivatives. 
 
 Experiment IV. Recall the formation of ozone. 
 
 O 2 + O O 3 32,400 cals. 
 
OXYGEN FAMILY. 177 
 
 SULPHUR, S. 
 
 At. Wt. 32.06 Mol. Wt, (S 2> S 4 , S e , S 8 ). 
 Properties oj Sulphur. 
 Experiment I. Transition Points. Plastic Sulphur. 
 
 Procure a few grams of roll sulphur; note its color, taste, odor, relative 
 density, structure, luster. Place 5-10 grams of the sulphur in a dry test 
 tube; heat gently, observing the yellow deposit of sulphur (flowers of sul- 
 phur) on the cooler portions of the tube as the experiment proceeds ; increase 
 the temperature of the tube until the sulphur melts. Determine the 
 melting point by means of the thermometer, then remove it; continue 
 to increase the supply of heat, noting the change in color of the fluid. 
 At 1 60 - 2 60 the yellow mobile fluid changes to a dark liquid so viscous 
 that the test tube may be inverted without spilling the sulphur. Heat to 
 boiling (448.5); observe that the fluid becomes less viscous. Collect a 
 portion of the sulphur vapor in a cold flask or bottle. 
 
 Pour the boiling sulphur from the test tube into a beaker of water. 
 Examine the cooled mass. Does it resemble roll sulphur ? Is it amorphous 
 or crystalline? It is known as plastic sulphur. Place a part of it on a 
 watch glass and set aside to be examined after a couple of days. Re- 
 sults ? Reserve the other portion for the following experiment. 
 
 Note. Plastic (elastic) sulphur and, in fact, all amorphous bodies 
 are spoken of as " supercooled liquids." See Lecture Notes. 
 
 Experiment II. Monoclinic Sulphur. 
 
 Melt 25-50 grams of sulphur in a porcelain evaporating dish or a 
 Hessian crucible; cool until a crust has barely formed over the surface; 
 perforate the crust with a glass rod, and pour off the liquid portion. 
 The sides and bottom of the dish will be covered with transparent, needle- 
 like crystals of sulphur. Set the monoclinic crystals aside for several 
 days, then observe that they have become opaque like roll sulphur. 
 
 Experiment III. Rhombic Sulphur. Solubility of the Allotropic Forms 
 of Sulphur. 
 
 (a) Place a small piece of roll sulphur in the bottom of a test tube; add 
 5 or 6 cm. 3 of CS 2 ; after shaking for some time, pour a cm. 3 of the liquid 
 upon a glass ;plate; the CS 2 will evaporate spontaneously depositing 
 crystals of rhombic (octahedral) sulphur. Examine the crystals with a 
 hand- magnifying glass or by the aid of a microscope. Are the crystals 
 soluble in H 2 O ? In CS 2 ? 
 
 (b) Repeat (a) using plastic amorphous sulphur, monoclinic (pris- 
 matic) sulphur, and flowers of sulphur. Test each form separately. 
 Record results. Sulphur crystallizes in how many forms? Such a 
 substance is given what name? 
 
178 EXPERIMENTAL CHEMISTRY. 
 
 Native sulphur is a yellow crystalline substance composed largely of 
 partially formed rhombic crystals. It has a sp. gr. of 2.06; melts at 
 114.5, an d boils at 448.5. Below 96, prismatic or monoclinic sulphur 
 is unstable and passes into the rhombic variety; above 96 rhombic 
 sulphur is unstable and passes into the monoclinic. The range of stable 
 rhombic sulphur extends from low temperature to 96 C.; that of mono- 
 clinic from 96 to 119.5, i ts melting-point. Amorphous sulphur, 
 likewise monoclinic, if allowed to stand at ordinary temperatures becomes 
 hard and brittle; it passes into the rhombic. 
 
 That which is of special interest is the fact that monoclinic sulphur 
 with a sp. gr. of 1.96 and a melting-point of 119.5, possesses properties 
 quite different from those possessed by the rhombic form. (The opaque- 
 ness which was attendant upon the transformation of the monoclinic 
 into rhombic crystals, is due to the fact that the rhombic crystals with 
 their greater sp. gr. (2.07) occupy less actual space than the monoclinic 
 or prismatic.) The difference between these two varieties of sulphur 
 suggests that which was observed between the two forms of oxygen, 
 namely, ordinary oxygen and ozone. (Monoclinic sulphur is frequently 
 referred to as the analogue of ozone.) It is quite doubtful, however, 
 whether an analogous explanation will satisfactorily meet the require- 
 ments of both cases. It seems quite certain in the case of sulphur that 
 the different energy contents of the two forms are in some way intimately 
 related to the difference in properties. The change from the monoclinic 
 to the rhombic form is accompanied by an evolution of a considerable 
 amount of heat. (Compare ozone.) It is evident that the molecule of 
 monoclinic sulphur contains more internal energy than the molecule of 
 rhombic sulphur. Favre and Silbermann (Ann. Chim. Phys.) treated 
 this, and found the following: 
 
 S(monoclinic) + O 2 SO 2 + 73,300 cal. 
 S (rhombic) + O 2 * SO 2 + 71,000 cal. 
 
 The difference, 2300 cal., is the thermal equivalent of the difference 
 between the energy contents of the two forms of sulphur. 
 
 Experiment IV. Union of Sulphur with Metals to Form Sulphides. 
 
 Place a small globule of mercury in a clean mortar containing a little 
 sulphur; rub the two substances together by use of pestle. What is the 
 color of the resulting powder? Compare it with the mercury sulphide 
 found on the end-shelf. 
 
 With the exception of platinum and gold, finely divided metals when 
 rubbed with sulphur combine with the latter to form sulphides. 
 
 Experiment V. Formation of Ferrous Sulphide. Hydrogen Sulphide. 
 
 To 5 grams of powdered sulphur in a mortar, add 7 grams of fine iron 
 filings; mix thoroughly and introduce mixture into a small test tube; heat 
 the tube until the contents glow; cool, and break the tube; examine the 
 
OXYGEN FAMILY. 179 
 
 fused mass. Is it a mixture or a compound ? To several small pieces in a 
 test tube, add dilute HC1 or H 2 SO 4 . Note the odor of the evolved gas, 
 hydrogen sulphide or, as it is sometimes styled, sulphuretted hydrogen. 
 Indicate the chemical reactions by means of equations. 
 
 H 2 -f S (rhombic) H 2 S + 2730 cals. 
 
 Experiment VI. Preparation and Properties of Hydrogen Sulphide. 
 Hydrosulphuric Acid. 
 
 Place several pieces of ferrous sulphide (iron sulphide) in the flask 
 of the apparatus used for generating hydrogen; insert the stopper; 
 through the thistle tube introduce 2 cm. 3 of strong HC1 or H 2 SO 4 . Ob- 
 serve the slow evolution of the hydrogen sulphide, H 2 S; add H 2 O slowly 
 until the evolution of gas becomes rapid. Explain action of H 2 O. 
 Collect some of the gas over water. Note its color, odor and its density 
 as compared with air. Inflame a bottle of the gas. Explain the deposit 
 on the inside of the test tube. 
 
 Determine the solubility of the gas in water. Test the action of an 
 aqueous solution of it upon litmus paper, and a solution of sodium car- 
 bonate. A water solution of H 2 S is frequently termed hydro sulphuric 
 acid or sulphrydic acid. 
 
 Raise the delivery tube from the water; attach a jet and ignite the 
 issuing gas if all the air has been expelled from the flask. Hold a cold 
 piece of porcelain for example, an evaporating dish or a mortar in the 
 darker portion of the flame. What substance, evidently in the free state 
 in the inner portion of the flame, is deposited upon the porcelain ? What 
 other elementary substance probably exists in a free state in the interior 
 of the flame? State the probable reason for the existence in the free 
 state of these two elementary substances. 
 
 Note. Use the remaining gas to saturate 50 cm. 3 of H 2 O in a flask. 
 This aqueous solution may be used for the following experiments if the 
 laboratory is not equipped with a generator which furnishes a constant 
 supply of H 2 S. 
 
 (H 2 S. Aq) = 4500 cal. 
 Experiment VII. Properties of Hydrogen Sulphide. Continued. 
 
 (a) Instability of hydrogen sulphide. In a clean test tube expose an 
 aqueous solution of the gas to the action of the air for several days. 
 Results ? Equation ? 
 
 (b) Reducing action of hydrogen sulphide. Pass a stream of H 2 S 
 through 5 cm. 3 of strong H 2 SO 4 - until there is evidence that SO 2 is being 
 evolved and free sulphur is deposited. Equation? Could H 2 SO 4 be 
 used to dry H 2 S ? 
 
 To 5 cm. 3 of a K 2 Cr 2 O 7 (potassium dichromate) solution add 2 cm.s 
 of HC1. What is the color of the solution? Pass H 2 S through the 
 solution until its color is green, due to the presence of CrCl 3 . Is free 
 
l8o EXPERIMENTAL CHEMISTRY. 
 
 sulphur deposited? Write the structural formula for K 2 Cr 2 O 7 . What 
 is the valence of the chromium in this salt? Structural formula for 
 CrCl 3 ? Valence of chromium in this salt ? Write equations showing 
 nature of the reactions. 
 
 The above may be repeated using KMnO 4 instead of K 2 Cr 2 O 7 . 
 
 To 3 cm. 3 of ferric chloride, FeCL 3 add a few drops of potassium 
 ferricyanide, K 3 Fe(CN) 6 . Note the color. Repeat, using ferrous 
 sulphate, FeSO 4 , instead of FeCl 3 . Note the color. Pass H 2 S through 
 separate solutions of FeCl 3 , and FeSO 4 , for two or three minutes; filter 
 and test each filtrate with K 3 Fe(CN) 6 . Explain results. Equations? 
 
 (c) Formation of sulphides. Allow hydrogen sulphide to come into 
 contact with a moist silver coin. Results ? Equation ? 
 
 (d) Determine the electrical conductivity of an aqueous solution of 
 H 2 S. What is your conclusion as to its degree of dissociation ? Is it a 
 strong or weak acid ? Why ? 
 
 Experiment VIII. Hydrogen Sulphide and Sulphides. 
 
 Introduce 3 cm. 3 of ZnSO 4 solution into a test tube add 5 cm. 3 of 
 water. Test the solution with litmus paper. Is the solution neutral? 
 Saturate solution with H 2 S. Does a precipitate of ZnS form? Is the 
 precipitate soluble in HC1 ? Test by adding 2 cm. 3 of said acid. Saturate 
 solution a second time with the H 2 S. Is a precipitate formed? Write 
 the equations for the above reactions showing the interaction of the ions. 
 
 Repeat the foregoing, using solutions of CuSO 4 , Cd(NO 3 ) 2 , NaCl, 
 MnSO 4 , BaCl 2 , H 3 AsO 3 , in place of ZnSO 4 . Classify the sulphides 
 according to their solubility in water, dilute acids and strong acids. 
 
 Sulphides of the alkalies and the alkaline earths are soluble in water; 
 all other metallic sulphides are insoluble. The behavior of the insoluble 
 sulphides towards acids may be summarized as follows: 
 
 (a) Sulphides decomposed by dilute acids like HC1 or H 2 SO 4 . Ex- 
 amples: FeS, ZnS, MnS. 
 
 (b) Sulphides not decomposed by dilute acids, but soluble in warm 
 strong HC1. Examples: PbS, Sb 2 S 3 , SnS, NiS, CoS. 
 
 (c) Sulphides not soluble in strong HC1, but are decomposed by aqua 
 regia or by a mixture of HC1 and KC1O 3 . Examples: HgS, As 2 S 3 . 
 
 Experiment IX. A Test for Sulphides. 
 
 (a) When soluble sulphides are decomposed by acids, HjS is evolved. 
 The presence of the gas may be detected (i) by its odor; (2) by its action 
 upon lead acetate. This latter test is usually made as follows: A 
 piece of filter paper moistened with lead acetate is held over the mouth 
 of the test tube in which the reaction occurs. The H 2 S causes a black 
 stain due to the formation of lead sulphide. 
 
 (b) Prepare a dilute solution of Na,jS or K 2 S. Note that the odor of 
 H 2 S is not present. Following (a), test for sulphides. Equations ? 
 
OXYGEN FAMILY. l8l 
 
 Experiment X. Optional. Polysulphides. Hydrogen Poly sulphide. 
 
 (a) To 20 cm. 3 of a solution of a soluble sulphide, e.g., Na^S, add a 
 pinch of pulverized sulphur; shake vigorously until the color of the 
 solution becomes reddish-yellow. After evaporation of the solution, 
 the substances remaining have a composition varying from Na^ to 
 Na 2 S 5 . Dissolve the residue in water. Divide the solution into two 
 parts. To one portion add strong HC1 and note the liberation of H 2 S 
 and free sulphur (rhombic). Add the other portion to 3 cm. 3 of strong 
 HC1. Observe that no H 2 S is evolved, but that a heavy yellow oil, 
 hydrogen pentasulphide (H 2 S 5 ) forms a distinct phase in the bottom of 
 the tube. H 2 S 5 is unstable in the presence of moisture and decomposes 
 into H 2 S and S. 
 
 (b) Add a solution of CaS 2 , (CaS x ) to HC1; the yellow oily liquid, 
 hydrogen polysulphide, separates. 
 
 Oxygen Derivatives of Sulphur. 
 
 Experiment XI. Formation of Sulphur Dioxide. 
 
 (a) Oxidation of sulphur by burning in the air. Set fire to a few 
 small pieces of sulphur upon the inverted cover of a porcelain crucible. 
 Observe the color of the flame and note the odor of the sulphur dioxide, 
 SO 2 , fumes. 
 
 (b) Roasting a pyrite, FeS 2 . Heat i gram of granular pyrite in a 
 hard glass test tube. Observe the deposit of sulphur on the cooler 
 portions of the tube, and notice the odor of evolved SO 2 . What is the 
 effect of heating those portions of the tube on which sulphur is deposited ? 
 
 (c) Reduction of H 2 SO 4 by sulphur. Heat a small piece of roll 
 sulphur with strong H 2 SO 4 in a test tube. Wave the escaping fumes 
 toward the nose. Odor? Equation? Pieces of charcoal may be substi- 
 tuted for the sulphur, and the above procedure followed. 
 
 (d) Reduction of H 2 SO 4 by copper. See Exp. XII. 
 
 S + O 2 SO 2 + 71,000 cal. 
 
 Would you infer from equation that SO 2 is a stable or an unstable 
 compound ? 
 
 Experiment XII. (Hood) Laboratory Preparation of Sulphur Dioxide. 
 
 Fit a 300 cm. 3 flask with a rubber stopper provided with thistle tube 
 and a delivery tube, and add two or three small bunches of copper 
 shavings or 10 grams of granulated copper; support the flask in ring 
 stand upon a wire gauze; connect flask (Fig. 29) with wash bottle, one- 
 third filled with strong H 2 SO 4 ; the wash bottle should be provided with 
 a doubly bent delivery tube, reaching almost to the desk-top. Add through 
 thistle tube 100 cm. 3 to 150 cm. 3 of strong H 2 SO 4 ; heat bottom of flask 
 gently until there is a brisk evolution of gas. Identify the gas. Fill three 
 bottles with the gas by downward displacement of air; stopper the bottles. 
 Saturate 25 cm. 3 of H 2 O with the gas; the delivery tube should be placed 
 
i8 2 
 
 EXPERIMENTAL CHEMISTRY. 
 
 i cm. above the surface of the water. What evidence is there that the 
 gas is dissolving? Proceed with the study of the properties of SO 2 . 
 Equation indicating action of strong H 2 SO 4 on Cu? 
 
 Experiment XIII. Properties of Sulphur Dioxide. 
 
 (a) Bleaching properties of sulphur dioxide. Place a moist strip 
 of colored calico and the moistened petals of a red flower in one bottle 
 of the gas. Set them aside until the close of the laboratory period. 
 Effect of gas? 
 
 FIG. 29. 
 
 (b) Using another bottle of the'gas, ascertain if SO 2 will burn or support 
 combustion. Conclusions ? 
 
 (c) Devise a method for determining the density of SO 2 as compared 
 with air. Record procedure and results. 
 
 (d) What is the effect of dry SO 2 upon dry blue litmus paper? 
 
 Experiment XIV. Optional. Liquid Sulphur Dioxide. 
 
 Note. The critical temperature of SO 2 is 156 C. The gas, SO 2 , 
 when cooled to the temperature of a freezing mixture of salt and ice, 
 condenses to a liquid which boils at about 8 C. 
 
 Pass a slow current of dry SO 2 through a long spinal tube, immersed 
 in a freezing mixture, into a U-tube, the arms of which are provided 
 with glass stopcocks. This latter tube, "sulphur dioxide condenser," 
 must also be placed in a freezing mixture. The liquid SO 2 may be 
 kept indefinitely in the closed tube. When gaseous SO 2 is required, 
 one of the stop cocks is opened. 
 
OXYGEN FAMILY. 183 
 
 Experiment XV. Aqueous Solution of Sulphur Dioxide. Sulphurous 
 Acid. 
 
 Test the action of the aqueous solution of SO 2 prepared in Exp. XII. 
 upon blue litmus paper. To i cm. 3 of the solution add 10 cm. 3 of H 2 O. 
 Taste a drop of the solution. Does the solution manifest the properties 
 of an acid, base or salt? Explain its formation. Equation? Heat 
 5 cm. 3 of the acid solution. Note the odor of the fumes. SO 2 is the 
 anhydride of what acid? What is the significance of the ous ending? 
 What is the termination of its corresponding salts ? Write the structural 
 formula of sulphurous acid. How many series of salts can H 2 SO 3 
 form ? What name is applied to such acids ? Would you say that 
 H 2 SO 3 is a very stable compound? Why? 
 
 SO 2 + H 2 O SO 2 , Aq + 7,700 cal. 
 
 Experiment XVI. Preparation of a Soluble Sulphite. Action of Strong 
 Acids Upon Soluble Sulphites. Test for Sulphites. 
 
 Place 10 cm. 3 of H 2 SO 3 in a test tube; add a solution of NaOH until 
 the odor of SO 2 has disappeared. Divide the sodium sulphite solution 
 into two parts. To one portion add a little HC1; to the other portion 
 add H 2 SO 4 . Is SO 2 evolved in both cases? Equations? 
 
 Experiment XVII. A Test for Sulphites. Reversible Reaction. 
 
 To 5 cm. 3 of a BaCl 2 solution add drop by drop several cm. 3 of H 2 SO 3 . 
 Observe that the white precipitate of BaSO 3 is not heavy. Write the 
 equation for the reaction. Is the precipitate soluble in HC1? What 
 acid was formed when BaSO 3 was precipitated? Will this acid have a 
 tendency to reverse the action of the H 2 SO 3 ? Is it probable that the 
 precipitation of BaSO 3 would be more complete if the acid were removed 
 as fast as it formed ? 
 
 Ascertain the effect of adding a solution of sodium or potassium 
 acetate to a solution of BaCl 3 to which has been added H 2 SO 3 . Is the 
 precipitation more complete ? Is the precipitate soluble in HC1? What 
 gas is evolved? Explain the further precipitation of BaSO 3 on the 
 addition of sodium acetate, NaC 2 H 3 O 2 , by use of " ionic" equations. 
 
 All sulphites except those of the alkali metals are insoluble in water. 
 Dilute acids decompose all sulphites with an evolution of SO 2 . The 
 larger number of the sulphites are converted into sulphides and sulphates 
 when heated. Those of the alkaline earths yield a metallic oxide and 
 S0 2 . 
 
 4 Na 2 SO 3 ->Na 2 S + 3Na,SO 4 , 
 BaSO 3 ->BaO + SO 2 . 
 
 Experiment XVIII. Reducing Action of Sulphurous Acid. 
 
 (a) What is the effect of adding H 2 SO 3 to a very dilute solution of 
 KMnO 4 acidulated with H 2 SO 4 ? Equations? 
 
184 EXPERIMENTAL CHEMISTRY. 
 
 (b) Substituting K 2 Cr 2 O 7 for KMnO 4 , repeat (a). Equations? 
 
 (c) Recall action of H 2 SO 3 on HIO 3 . Was reduction accompanied 
 by oxidation ? Explain. 
 
 Sulphurous acid in the presence of more powerful reducing agents 
 may act as an oxidizing substance as follows: 
 
 Experiment XIX. Oxidizing Action of Sulphurous Acid. 
 
 (a) Add 2 cm. 3 of HC1 to 5 cm.3 of SnCl 2 ; add 2 cm.3 of H 2 SO 3 ; warm 
 gently; observe the formation of a yellow precipitate of SnS 2 . As the 
 reactions usually proceed rather slowly, allow the test tube and contents 
 to remain undisturbed for 10 minutes to 15 minutes. The H 2 SO 3 is 
 reduced to H 2 S, and the SnCl 2 is oxidized to SnCl 4 . 
 
 (1) 3SnCl 2 + 6HC1 + H 2 SO 3 3SnCl 4 + 3H 2 O + H^Si 
 
 (2) SnCl 4 + 2H 2 S SnS 2 + 4HC1. 
 
 (b) Nascent hydrogen reduces sulphurous acid. Place a piece of 
 granulated zinc in the bottom of a test tube; add a little dilute HC1; add a 
 few drops of H 2 SO 3 . Test for the presence of H 2 S by holding a piece of 
 filter paper moistened with a drop of lead acetate over the mouth of 
 the tube (?). Equations? 
 
 Experiment XX. Oxidation of Sulphurous Acid. Sulphuric Acid. 
 
 (a) To 3 cm.3 o f H 2 SO 3 add a little H 2 O 2 ; shake, then add a small 
 quantity of a BaCl 2 solution. Is the precipitate which forms soluble in 
 HC1? Recall the solubility of BaSO 3 in HC1. The white precipitate 
 insoluble in HC1 is BaSO 4 . What acid was evidently present in the 
 solution before the BaCl 2 was introduced? Explain its formation. 
 Equations? Using H 2 SO 4 from the reagent shelf, test its action on 
 BaCl 2 . Is the precipitate soluble in HC1? How can you distinguish 
 between a sulphite and a sulphate ? 
 
 (b) Repeat (a) using "bromine water." Results? Equations? 
 
 (c) Slow oxidation of H 2 SO 3 by oxygen of the air. Expose 10 cm. 3 
 of H 2 SO 3 to the action of the air for several hours. Test it with BaCl 2 . 
 Was BaSO 3 or BaSO 4 formed ? Give reasons for answer. Equations ? 
 
 (b) Rapid oxidation of H 2 SO 3 by oxygen of the air by use of a catalyser, 
 N 2 O 3 . Prepare the gaseous catalyser, N 2 O 3 , as follows: Put 10 or 
 15 grams of granulated copper or a bunch of copper shavings in a 
 generating flask; add sufficient water through the thistle- tube to cover 
 the lower end of the latter; then introduce 20 cm. 3 of concentrated HNO 3 . 
 Observe the formation of a reddish-brown gas which is a mixture of N 2 O 3 
 and N 2 O 4 . The colorless gas, NO, is the initial product formed by the 
 action of HNO 3 on copper, but in the presence of air, it quickly oxidizes 
 to the above-mentioned gases. Fill a small flask with NO by displace- 
 ment of water; allow air to enter flask that N 2 O 3 may be formed, then 
 pour 2 cm. 3 of H 2 SO 3 into the flask; fit a stopper in place and shake. 
 
OXYGEN FAMILY. 
 
 Observe that the color disappears. Remove stopper until gas is again 
 colored; insert stopper and shake (?). Repeat this operation several 
 times, then by use of BaCl 2 and HC1 prove the presence of H 2 SO 4 . 
 
 Experiment XXI. (L. T.) 
 Chamber Process. 
 
 The Manufacture of Sulphuric Acid. 
 
 The purpose of the following experiment is to acquaint the student 
 with the "chemistry" of the sulphuric acid industry. References and 
 lectures should supplement the experiment. A flask or glass globe of two 
 or three liters' capacity is fitted with a cork through which passes five de- 
 livery tubes reaching to the center of the flask (Fig. 30). The middle tube 
 e serves as an escape for the excess of gases introduced through the other 
 
 FIG. 30. 
 
 tubes. Through a tube d air can be forced into the flask (chamber) 
 by a bellows or a bicycle pump. The other three tubes are connected 
 by glass tubing with separate generating flasks a, b and c. a is the flask 
 in which SO 2 is prepared by heating copper shavings and concentrated 
 H 2 SO 4 . (The SO 2 may be prepared by heating granular pyrite in a 
 boat placed in hard glass tubing. A current of air should be drawn 
 through the tube over the pyrite by use of an aspirating bottle.) Nitric 
 oxide, N 2 O 2 , which in the presence of oxygen oxidizes to N 2 O 3 , the cataly- 
 ser, is generated in b by the action of dilute HNO 3 on copper shavings. 
 (Nitric acid vapor may be used instead.) Steam, H 2 O, is supplied by 
 boiling water in flask c. If the products (gases) from a and b together 
 with air are introduced without a sufficient supply of steam, a frost-like 
 
1 86 EXPERIMENTAL CHEMISTRY. 
 
 deposition of nitrosyl-sulphuric acid crystals is observed upon the walls 
 of the flask. 
 
 H O H. 
 
 \ / 
 
 O + 2SO 2 + N 2 O 3 + O 2 > 2SO 2 
 
 / \ 
 
 H O NO. 
 
 Now increase the supply of steam and note the disappearance of the 
 crystals and the formation of a heavy oily fluid: 
 
 O H OH 
 
 / / 
 
 2 S0 2 + H 2 * 2 S0 2 + N 2 3 . 
 
 \ \ 
 
 O NO OH 
 
 When 10 cm. 3 to 15 cm. 3 of liquid have collected in the flask, the 
 experiment may be interrupted. Place the liquid in a test tube and 
 reserve it for following experiments. 
 
 Experiment XXII. Properties of Sulphuric Acid (Hydrogen Sulphate). 
 
 (a) Dehydrating action of sulphuric acid. Action of sulphuric acid 
 with organic matter. Remove the head of a match; dip the \vood into 
 hot concentrated H 2 SO 4 . Results? Write your name with dilute 
 H 2 SO 4 upon a piece of paper. Dry the paper. Results? 
 
 To a little sugar in the bottom of a test tube add strong H 2 SO 4 ; heat 
 gently. Results? Each of the above substances is composed largely 
 of C, H, and O. What is the dark-colored substance which remains in 
 each case ? 
 
 Repeat each of the above experiments with the oily fluid prepared in 
 Exp. XXI. Make a record of results. Are its properties identical 
 with the shelf-reagent, H 2 SO 4 ? Is sulphuric acid a efficient drying agent ? 
 
 (b) To i cm. 3 of the fluid prepared in Exp. XXI, add 5 cm. 3 of water. 
 Test with blue litmus paper. Repeat above, using H 2 SO 4 . Results? 
 
 (c) Try the action of the "acid" on a piece of zinc in a test tube. 
 What gas is evolved? Repeat with H 2 SO 4 . Compare results. 
 
 (d) Add BaCl 2 to a portion of the acid prepared in above experiment. 
 Is the precipitate soluble in HC1? 
 
 Does H 2 SO 4 yield similar results? Drawing your inference from 
 the results of the above experiments, identify the liquid prepared in 
 Exp. XXI. Write the structural formula for H 2 SO 4 . 
 
 Experiment XXIII. Hydration of Sulphuric Acid. 
 
 (a) Into 5 cm. 3 of water in a test tube add 10 cm. 3 of strong H 2 SO 4 . 
 Is there a change in the temperature of the liquids ? What is the probable 
 explanation of the thermal phenomenon? 
 
 H 2 SO 4 + 2H 2 O H 2 SO 4 , Aq + 17,850 cal. 
 
 Is the solution more or less stable than the pure acid ? Why ? 
 
OXYGEN FAMILY. 1 87 
 
 (b) Optional. The following interesting experiment is suggested by 
 Freer: Add 9.8 grams of concentrated H 2 SO 4 to 1.8 grams of water; place 
 the liquid in a small flask and surround the latter with a freezing mixture 
 of ice and salt; crystals of H 2 SO 4 .H 2 O (H 4 SO 5 ) will form; warm the 
 crystals until they are melted (80 C.), then add 1.8 grams of water; again 
 place in the freezing mixture and crystals of H 2 SO 4 .2H 2 O (H 6 SO 6 ) will 
 separate. Further addition of water is not accompanied by increase 
 of temperature. 
 
 H 2 SO 4 , Normal Sulphuric Acid. 
 
 /OH O x / O H 
 H 2 S0 4 , S0 2 (OH) 2 , S0 2 ( >S( 
 
 X OH, O/ X H. 
 
 H 2 SO 4 .H 2 O, Tetra hydroxyl Sulphuric Acid. 
 
 /OH /OH 
 
 "/OH //OH 
 
 H 4 S0 5 , SO/ , O = S/ , 
 
 X OH X OH 
 
 X OH X OH 
 
 H 2 SO 4 .2H 2 O, Hexahydroxyl Sulphuric Acid. 
 
 /OH 
 /OH 
 
 /OH 
 H 6 S0 6 , S(OH) 6 , S< 
 
 X OH 
 X OH 
 X OH 
 
 Experiment XXIV. Dissociation of Sulphuric Acid. 
 
 Pour 4 cm. 3 of concentrated H 2 SO 4 into a clean dry test tube; clamp 
 the tube to the ring stand; suspend a 2oo-3oo C. thermometer so that 
 the bulb will be immersed in the acid; heat the test tube gently with a 
 small flame. Take the reading on the thermometer at the first appear- 
 ance of the heavy, dense white fumes of SO 3 . Caution. Hot H 2 SO 4 
 produces severe burns. Note the thermometer reading frequently, 
 that the thread of mercury may not be allowed to approach too near 
 the top of the tube. Write the equation indicating the dissociation of 
 H 2 SO 4 . What is the difference between dissociation and decomposition ? 
 
 Experiment XXV. Reduction of Sulphuric Acid. 
 
 Recall Exp. XI. (r), (d). Equations? 
 
 Experiment XXVI. Soluble and Insoluble Sulphates. 
 
 Test the action of a solution of Na 2 SO 4 or dilute H 2 SO 4 upon separate 
 solutions of each of the following substances: AgNO 3 , NaCl, CaCl 2 , 
 
1 88 EXPERIMENTAL CHEMISTRY. 
 
 Pb(N0 3 ) 2 , Cu(N0 3 ) 2 , KN0 3 , BaCl 2 , Sr(NO 3 ); NH 4 C1, FeCl 3 . Write 
 the ionic equations for each reaction. Underscore the formula repre- 
 senting the precipitate in each case. 
 
 Experiment XXVII. Optional. Decomposition of Isoluble Sulphates. 
 
 Insoluble sulphates, e.g., BaSO 4 , may be decomposed by fusion with 
 Na^CC^; Na^SC^, which is soluble in water, being formed. The fused 
 mass is extracted with water and the solution made slightly acid with 
 HC1, after which BaCl 2 is added. The substances are usually fused in a 
 crucible. 
 
 Note. The above method is quite generally used for converting 
 insoluble salts into corresponding salts of the alkali metals. Practically 
 all of the salts of sodium and potassium are soluble in water, hence the 
 use of Na-jCOg or K 2 CO 3 as a flux. 
 
 Most sulphates are soluble in water. Silver, calcium, strontium, 
 lead and barium are but slightly soluble in water. Their solubility 
 decreases in the order named. 
 
 Experiment XXVIII. Optional. Reduction of Sulphates and Sulphites. 
 
 (a) Mix a small quantity of a sulphate or a sulphite with two or three 
 times its weight of pure anhydrous Na^COg; place a portion of the 
 mixture in a depression in a piece of charcoal; using a blow-pipe, direct 
 the reducing flame upon the mixture. Place the fluid mass upon a silver 
 coin or a piece of filter paper moistened with lead acetate; allow a drop 
 of dilute HC1 to fall upon the mass. Explain the formation of a black 
 stain. Equations ? 
 
 Experiment XXIX. Optional. Sulphur Trioxide. The Manufacture 
 of Sulphuric Acid by the "Contact Process." 
 
 (a) Put 15 cm. 3 of fuming sulphuric acid (disulphuric or pyrosul- 
 phuric acid, H 2 S 2 O 7 ) in a small dry retort provided with a glass stopper 
 and connected with a dry receiver immersed in water. Keep the re- 
 ceiver cool. Heat the flask gently, and observe evolution of the heavy 
 SO 3 fumes which condense to a solid in the receiver. Remove some of 
 the solid by means of a glass rod and put it in a little water. Is heat 
 evolved ? Place other portions on pieces of paper and wood. Results ? 
 
 (b) Sulphur dioxide and oxygen do not combine rapidly under or- 
 dinary conditions to form sulphur trioxide. At higher temperatures 
 the reaction proceeds relatively slowly. Above 400 C. , S O 3 is decomposed 
 into SO 2 and O. It has been found that if the two gases are passed 
 simultaneously over heated finely-divided platinum they combine to form 
 SO 3 which combines readily with water to form H 2 SO 4 . The above is a 
 general statement of the principles involved in the manufacture of H 2 SO 4 
 by the "contact process." 
 
 Prepare platinized asbestos as previously directed; place it in a long 
 
OXYGEN FAMILY. 189 
 
 hard glass tube. Devise a method for passing oxygen and sulphur 
 dioxide through the tube. The gases should be dry. Collect the gases 
 in a dry receiving flask. Is there any evidence of SO 3 being formed? 
 Now gently heat the tube beneath the asbestos. Results? When the 
 products are added to water, the solution may be tested for sulphuric 
 acid. 
 
 50 2 + O > SO 3 + 32,100 cal. 
 
 S '" + O 3 SO 3 (liquid) + 103,200 cal. 
 
 50 3 + H 2 O H 2 SO 4 + 21,300 cal. 
 (H 2 SO 4 ,Aq) = 17,800 cal. 
 
 (S, t> 8 , Aq) - ? cal. 
 (H 2 ,S,0 4 ) = ? cal. 
 
 Experiment XXX. Preparation of a Soluble Thiosulphate. 
 
 Dissolve 5 grams of sodium sulphite, NaSO 3 in 20 cm. 3 of H 2 O, then 
 add 3 grams of finely divided sulphur, e.g., flowers of sulphur. Boil 
 the mixture for 10 or 15 minutes. Replace the water which has evapor- 
 ated ; filter. Pour the filtrate into a small beaker or a crystallizing dish. 
 Sodium thiosulphate, Na^Og^ELjO, commonly called hyposulphite 
 of soda ("hypo"), will crystallize out of the solution. Acids, e.g., 
 H 2 SO 4 , decompose thiosulphates with the formation of a sulphate, 
 sulphurous acid and the liberation of free sulphur. Try the action of 
 acids upon the crystals. Results ? Equations ? 
 
 Experiment XXXI. Preparation of an Insoluble Thiosulphate. 
 Decomposition of Thiosulphates by Heat. 
 
 (a) To a few cm. 3 of a solution of a soluble thiosulphate add a few 
 drops of a AgNO 3 solution. Observe the color of the precipitate. Heat 
 the mixture and note the formation of black-silver sulphide. Is SO 2 
 evolved ? Is free sulphur liberated ? Equations ? 
 
 (6) Lead acetate may be substituted for AgNO 3 and (a) repeated. 
 
 Experiment XXXII. Oxidation of Sodium Thiosulphate to a Tetra- 
 thionate by the Action of Iodine. 
 
 Prepare a solution of sodium thiosulphate; add a little starch water. 
 Dissolve a small quantity of iodine in a potassium iodide solution. To 
 5 cm. 3 of the thiosulphate solution add the solution of iodine drop by 
 drop until the solution turns blue. Explain. Equations? 
 
 Note. The titration of iodine against sodium thiosulphate with starch 
 as an indicator is a process frequently used in analytical chemistry. 
 
 Experiment XXXIII. (Quant.) Compare the electrical conductivi- 
 ties of pure concentrated H 2 SO 4 and a 5N solution of same acid. Ex- 
 plain results. The activity of an acid depends on what? Would you 
 
1 90 EXPERIMENTAL CHEMISTRY. 
 
 infer that H 2 SO 4 in aqueous solution is active as an acid? Are the elec- 
 trical conductivities of acids related in any way to their activities ? Ex- 
 plain. Write the ionic equation indicating the different particles present 
 in an aqueous solution of sulphuric acid. 
 
 Experiment XXXIV. (Quant.) Specific Gravity of Sulphuric Acid. 
 
 Devise a method for determining the specific gravity of sulphuric 
 acid. What is the sp. gr. of the concentrated H 2 SO 4 which is found 
 among the shelf-reagents ? 
 
 Experiment XXXV. (Quant.) Contraction in Volume When Sul- 
 phuric Acid and Water are Mixed. 
 
 By means of a pipette place 10-20 cm. 3 of concentrated H 2 SO 4 in a 
 long narrow test tube or a colorimetric tube; carefully add one-half its 
 volume of water; mark the volume occupied by the two liquids. Now 
 mix the two liquids by careful manipulation of the tube. Note the 
 change in temperature of the liquids. Cool the mixture and determine 
 the amount of contraction. Does the acid undergo hydration ? 
 
 Although several factors may contribute to the contraction, it is known 
 that " water of hydration" occupies less space than ordinary water. 
 
 OXACIDS OF SULPHUR. 
 
 H 2 SO 3 Sulphurous acid. 
 
 H 2 SO 4 Sulphuric acid. 
 
 H 2 S 2 O 3 Thiosulphuric acid. 
 
 H 2 S 2 O 4 Hyposulphurous acid. 
 
 H 2 S 2 O 5 Pyrosulphurous acid. 
 
 H 2 S 2 O 6 Dithionic acid. 
 
 H 2 S 2 O 7 Pyrosulphuric acid. 
 
 H 2 S 2 O 8 Persulphuric acid. 
 
 H 2 S 3 O 6 Trithonic acid. 
 
 H 2 S 4 O 6 Tetrathionic acid. 
 
 H 2 S 5 O 6 Pentathionic acid. 
 
 H 2 S 6 O 6 Hexathionic acid. 
 
 SELENIUM, SC. 
 
 At. Wt. 79.2 Mol. Wt. 158.4. 
 See Text-book and Lecture Notes. 
 
 TELLURIUM, T6. 
 
 At. Wt. 127.6. Mol. Wt. (?). 
 See Text-book and Lecture Notes. 
 
OXYGEN FAMILY. 
 
 The following table gives a general view of the similarity of the ele- 
 ments of this group: 
 
 Oxygen Sulphur Selenium Tellurium 
 16.0 32.06 79.2 127.6 
 Gas Solid Solid Solid 
 Colorless, Pale- Grayish- Silver- 
 Yellow White White 
 
 Physical Properties 
 Atomic weight 
 State or phase 
 Color 
 
 Specific gravity 
 
 I-I3 
 
 2.02 
 
 4-5 
 
 6.3? 
 
 
 (at 182) 
 
 (Rhombic) 
 
 (Crystalline) 
 
 
 Specific heat 
 
 
 
 
 s 
 
 
 
 
 ^ 
 
 Melting-point 
 
 Below 223 111.5 
 
 217 
 
 452 
 
 Boiling-point 
 
 -182 
 
 445 
 
 675 
 
 1400 
 
 Chemical Properties 
 
 
 
 
 
 Heat of formation 
 
 H 2 O 
 
 H 2 S 
 
 H 2 Se 
 
 H 2 Te 
 
 of H- derivatives 57 
 
 ,000 cal. 2, 
 
 730 cal. 
 
 -5,400 cal. 
 
 19,400 cal. 
 
 Stability of H- 
 
 
 
 
 
 HATTVQ fi"\7 r p k c 
 
 ^s 
 
 
 
 <^ 
 
 LLCllVctllVCo 
 
 Heat of formation 
 
 2 ,0(gas) 
 
 S0 2 (gas) 
 
 SeO 2 (solid) 
 
 TeO 2 (solid) 
 
 of O- derivatives 32,400 cal. 
 
 71,000 cal. 
 
 57,700 cal. 
 
 77,000 cal. 
 
 Oxacids 
 
 
 H 2 SO 3 
 
 H 2 S 3 3 
 
 H 2 Te0 3 
 
 
 
 ? 
 
 (solid) 
 
 (solid) 
 
 
 
 H 2 SO 4 
 
 H 2 Se0 4 
 
 H 2 TeO 4 
 
 
 
 (liquid) 
 
 (liquid) 
 
 (liquid) 
 
 Heat of formation 
 
 0,CL 
 
 S 2 C1 2 
 
 Se 2 Cl 2 
 
 TeCl 4 
 
 of Cl-derivatives 18,000 cal. 
 
 14,300 cal. 
 
 22,200 cal. 
 
 77,400 cal. 
 
 Metallic Properties 
 
 
 
 
 
 
 
 
 _ 
 
 (Physical) 
 
 
 
 
 
 PROBLEMS. 
 
 1. How many grams of sulphur in il. of H 2 SO 4 , sp. gr. 1.84? How 
 many grams of oxygen? Of hydrogen? 
 
 2. How many grams of H 2 SO 4 will be required to neutralize 10 grams 
 ofBa(OH) 2 ? 
 
 3. How many cm. 3 of a 5N solution of H 2 SO 4 will be required to 
 neutralize 50 cm. 3 of iN of NaOH? 
 
 4. Calculate the heat of formation of anhydrous sulphuric acid from 
 its elements, i.e., (S, O 4 , H 2 ) = ? Ans. 192,900 cal. 
 
 5. What is the heat of formation of sulphuric acid from its elements 
 in dilute aqueous solution, i.e. (H 2 , S, O 4 , Aq) = ? Ans. 210,700 cal. 
 
CHAPTER XVIII. 
 
 NITROGEN AND THE ATMOSPHERE. THE HELIUM FAMILY- 
 NITROGEN, N. 
 At. Wt. 14.01. Mol. Wt. 28.02. 
 
 Nitrogen is a colorless, odorless and tasteless gas. At ordinary tem- 
 peratures it is chemically inert. 
 
 Preparation and Properties of Nitrogen. 
 Experiment I. Preparation of Nitrogen. 
 
 (a) Assemble the parts of the oxygen or hydrogen generator. Place 
 about 5 grams of sodium or potassium nitrite, NaNO 2 , KNO 2 , and 2 grams 
 of ammonium chloride, NH 4 C1, in the generating test tube or flask; add 
 10-15 cm. ^ of water; insert stopper provided with delivery tube; clamp 
 the tube to ring-stand; heat gently by waving the flame under the tube. 
 Be careful not to overheat the mixture. Fill three bottles with the gas by 
 displacement of water. Proceed with Exp. II. 
 
 (1) NaN0 2 + NH 4 C1 -> NH 4 NO, + NaCl 
 
 (2) NH 4 N0 2 N 2 + 2H,0. 
 
 (b) Alternative method. The gas may be prepared by heating ammo- 
 nium nitrite, NH 4 NO 2 , which decomposes into water and nitrogen. 
 
 Experiment II. Properties of Nitrogen. 
 
 Has the gas color, odor, taste? Is it inflammable? Will it support 
 combustion? Does it unite readily with other elements? Prove its 
 inertness in this respect. Is the gas heavier or lighter than air? 
 
 Experiment III. (Quant.) Determination of the Weight of a Liter of 
 Nitrogen. 
 
 Provide a round-bottomed 250 cm.s flask (Fig. 31) with a rubber cork 
 through which passes a piece of glass tubing about 8 cm. in length; attach a 
 piece of rubber tubing about 5 or 6 cm. long to the outer end of the glass tube 
 and wire it firmly. Pour about 30 cm. 3 of water into the flask, then firmly 
 press the stopper into place; make a mark on the neck of the flask at the 
 bottom of the stopper so as to be able to determine the exact contents of 
 the flask when the stopper is in place. Boil the water with a small flame 
 until all the air has been expelled from the flask. Allow the steam to 
 escape for 5 or 6 min., then close the rubber tube with a strong clip, and 
 
 192 
 
NITROGEN AND THE ATMOSPHERE. THE HELIUM FAMILY. 193 
 
 quickly remove the flame. Attach a fine wire to the neck of the flask by 
 which it may be suspended during the process of weighing. When the 
 flask has cooled to the temperature of the laboratory, wipe and carefully 
 weigh it. Read the temperature and barometric pressure in the balance- 
 room. Place a beaker of water so that the rubber siphon tube of an aspira- 
 tor bottle (Fig. 32) partially filled with nitrogen is made to dip beneath the 
 surface of the water. Connect the flask with the delivery tube of the 
 aspirator; the clip is gradually opened allowing a slow stream of nitrogen 
 to enter the flask. Now raise the beaker of water so that the water in it 
 will be at a higher level than that in the aspirator. Close the clip, dls- 
 
 FIG. 31. 
 
 FIG. 32. 
 
 connect from the aspirator delivery tube. Open the clip for an instant to 
 establish atmospheric pressure in the flask, then weigh again. The in- 
 crease in weight represents the weight of the nitrogen. From the table 
 in the appandix ascertain the aqueous tension at the observed temperature; 
 subtract this from the barometric pressure. Determine the volume of the 
 flask by filling with water to the mark placed on the neck and weighing it. 
 Calculate the weight of a liter of nitrogen at o C. and 760 mm. What is 
 the ratio of the weights of equal volumes of hydrogen and nitrogen? 
 What is the weight of the molar volume of the gas ? 
 
 THE ATMOSPHERE. 
 
 The atmosphere, or air as it is usually termed, is the gaseous, envelope 
 surrounding the earth. It is a mixture, the components of which may 
 be divided into two classes: (a) those components, oxygen, nitrogen, argon 
 and other members of the helium group, which are practically constant 
 in amount; (b) components which are variable in amount, e.g., carbon 
 dioxide, water vapor, ozone, dust, traces of ammonia, minute quantities 
 of solids like ammonium nitrate and ammonium carbonate, organic mat- 
 ter, etc. 
 
 The comparatively recent discovery of five gaseous elements in the air 
 has given a new impetus to the investigation of the composition of the 
 atmosphere. It is rather remarkable that these five elements constitute 
 a natural group, known as the helium group. The student is referred to 
 the well-known work by Sir William Ramsay, "The Gases of the Atmos- 
 phere." 
 
 Components of the Atmosphere. 
 13 
 
I 94 
 
 EXPERIMENTAL CHEMISTRY. 
 
 Experiment IV. (Quant.) Determination of the Amount of Oxygen 
 in the Air. Nitrogen. 
 
 (a) Absorption or Pyrogallate Method. See Exp. X. " Oxygen." 
 Repeat experiment. Calculate the percentage of oxygen [by volume 
 in the air. Test the residual gas as to odor, taste, inflammability. What 
 
 is the color of the gas ? Will it support combustion ? 
 Are its properties similar to those of nitrogen ? If this 
 residual gas were pure nitrogen instead of nitrogen 
 mixed with traces of other gases, what would be the per- 
 centage of nitrogen in the air ? 
 
 Note. HempeFs burette and compound pipette as 
 
 fc , .,/ IG "J% N apparata in which to measure and absorb the gas may 
 
 (Smith and Keller.} , rj , . , - Al L , , T ,/ 
 
 be used instead of that suggested above. See Hempels 
 
 "Gas Analysis," translated by L. M. Dennis. 
 
 (b) Explosion method. Substitute Hempel's explosion pipette (Fig. 33) 
 for the absorption pipette used in (a), and proceed as follows: To 25 cm. 3 
 of air contained in the burette add 25 cm. 3 of pure hydrogen, then pass the 
 mixture into the explosion pipette; close the pipette. (Instructions.) Pass 
 an electric spark through the mixture. Return the residual gas to the 
 burette and measure. What is the composition of this gas ? How much 
 of the contraction was due to oxygen ? Calculate the percentage of oxygen 
 in the air. 
 
 Experiment V. Removal of Oxygen from the Air. Determination of 
 the Approximate Percentage of Nitrogen in the Air. 
 
 (a) Place a small piece of yellow phosphorus either in an evaporating 
 dish floating on the water or in a small crucible resting in a depression 
 
 FIG. 34. 
 
 in a large cork which floats on the water in the pneumatic trough. Ignite 
 the phosphorus and quickly place a bell-jar or an inverted beaker over 
 the crucible (Fig. 34). Keep the rim below the surface of the water that no 
 gas may escape. Hold the bell-jar in this position until the dense heavy 
 
NITROGEN AND THE ATMOSPHERE. THE HELIUM FAMILY. 195 
 
 fumes of phosphorus pentoxide, P 2 O 5 , disappear. Note that the water has 
 risen in the jar. Lower the bell-jar until the water on the inside and out- 
 side have the same level; mark this level on the jar. Place a glass plate 
 over the mouth of the bell- jar and quickly invert it; test the gas with a 
 burning match. Results? Test the water in the trough with blue lit- 
 mus paper. Explain. How many cm. 3 of water are required to fill the 
 jar to the mark placed in it? How many cm. 3 are re- 
 quired to fill it? Calculate the percentage of nitrogen in 
 the air. Of Oxygen. Equations? 
 
 (b) If a piece of yellow phosphorus is inserted into a 
 measured volume of air contained in a graduated tube 
 (Fig. 35) over water or mercury, the phosphorus will com- 
 bine slowly with the air forming P 2 O 5 which unites with 
 water to form phosphoric acid. The phosphorus should be 
 allowed to remain in contact with the enclosed air for 24 
 to 48 hours. The water will slowly rise in tube to take 
 the place of oxygen. In determining the volume of the 
 residual gas the usual corrections should be applied. The 
 gas may be tested for nitrogen. 
 
 Experiment VI. Preparation of "Atmospheric Nitro- 
 gen." 
 
 Fit each end of a piece of hard glass tubing 25-30 cm. 
 long, with a cork through which passes a short piece of 
 glass tubing; clamp the tube in a horizontal position; * IG ' 35- 
 place a large plug of copper turnings in the middle of the 
 tube; connect tube in series with wash bottles, No. i and No. 2, con- 
 taining respectively a strong solution of KOH and concentrated H 2 SO 4 ; 
 the latter wash bottle is connected with an aspirating bottle filled with 
 mercury or water. Heat the copper red-hot; open the stop cock on the 
 siphon of the aspirator and allow the liquid to be siphoned off drop by 
 drop or in a slow stream. When the system has been swept free of air, 
 quickly substitute another aspirator which has been previously filled. 
 When three-fourths of the liquid has run out, close stop cock on siphon and 
 disconnect the aspirator from the wash bottle and attach a delivery tube; 
 raise the end of the siphon and attach a small funnel; pour water into 
 the funnel; open stop cock on siphon; the water will drive the gas out 
 through the delivery tube. Collect several bottles of the gas by displace- 
 ment of water, and apply the usual tests for nitrogen. Examine the 
 copper. Explain change of color. 
 
 Experiment VII. Presence of Water Vapor in the Atmosphere. 
 
 (a) Place small quantities of calcium chloride, CaCl 2 , and phosphorus 
 pentoxide, P 2 O 5 , on separate watch-glasses and expose them to the air. 
 Examine them after 10 or 15 min. Results? Conclusions? Touch the 
 moist P^,O 5 with a piece of blue litmus paper. Results? Explain. 
 Equation ? 
 
196 EXPERIMENTAL CHEMISTRY. 
 
 Which of the two substances apparently absorbs moisture the more rap- 
 idly ? Name another familiar substance which will remove moisture from 
 the air. 
 
 (b) A quantitative determination of the amount of moisture in the air 
 may be made by drawing a given volume of air through a U-tube filled 
 with glass beads and P 2 O 5 by means of an aspirator. The increase in 
 weight of the P 2 O 5 will represent the amount of moisture absorbed from a 
 known volume of air. (Instructions.) Substances which absorb moist- 
 ure are said to be hygroscopic. 
 
 Experiment VIII. Presence of Carbon Dioxide in the Atmosphere. 
 
 (a) Place 5 cm. 3 of a clear solution of calcium hydroxide, Ca(OH) 2 , 
 or barium hydroxide, Ba(OH) 2 , in a test tube. By use of a glass tube 
 force air (carbon dioxide) from the lungs through the solution until the 
 latter has a milk-like appearance, due to the formation of calcium car- 
 bonate, CaCO 3 . Equation? 
 
 Half fill another test tube or a wash bottle with a Ca(OH) 2 solution 
 (lime-water) and draw air through the clear solution for several minutes by 
 means of a filter pump or an aspirator. (If all the apparatus is not avail- 
 able expose the solution to the air.) Results? Conclusions as to the 
 presence of CO 2 in the air? 
 
 (b) Repeat the above experiment, but pass the air through a strong 
 solution of NaOH or KOH before it is allowed to enter the Ca(OH) 2 
 solution. Results ? Explain. 
 
 (c) The student is referred to works on "gas analysis" for methods of 
 making a quantitative determination of the CO 2 in the air. Report one 
 method. 
 
 Experiment IX. Absorption of Oxygen and Nitrogen by Magnesium 
 and Calcium. Argon, etc. 
 
 Prepare a mixture composed of equal parts of magnesium powder and 
 freshly-ignited calcium oxide. Keep the mixture perfectly dry. Intro- 
 duce a few grams of this mixture into a piece of combustion tubing sealed 
 at one end and about 20 cm. in length; clamp the tube in a horizontal 
 position; connect the open end of the tube with a delivery tube of smaller 
 bore and about 80 cm. long. The delivery tube should be bent so that 
 its free end may be immersed in a vessel of mercury 76 cm. below the 
 combustion tubing. Heat the mixture gently, then gradually increase the 
 temperature to that of a powerful burner. Continue to heat the tube 
 until the mercury no longer rises in the vertical tube. What is the resid- 
 ual gas ? What is the probable composition of the solid matter in the 
 tube? 
 
 Experiment X. Table of the Composition of the Air. 
 
 Prepare a tabulated statement of the average composition of the at- 
 mosphere. Before attempting to prepare the table, consult various ref- 
 erence w r orks. Give the names and authors of the books consulted. 
 
NITROGEN AND THE ATMOSPHERE. THE HELIUM FAMILY. 197 
 
 Experiment XI. (Quant.) Determination of the Weight of a Liter of Air. 
 
 The weight of a liter of air may be determined by a method similar 
 to that described under " Nitrogen." 
 
 THE HELIUM FAMILY. 
 
 The members of this natural group exhibit many close resemblan- 
 ces, e.g., remarkable inertness, monatomic nature of their molecules, 
 etc. They also manifest that gradation of properties which has been 
 observed in the other natural groups of elements, e.g., atomic volumes, 
 densities, refractive indices, etc. At first there seemed to be no place 
 within the periodic system for these elements, but it has been pointed 
 out that they may not unnaturally occupy the positions of transition 
 elements between the highly negative elements, the halogens, and the 
 strongly positive elements, the alkali metals. With the exception of argon 
 whose atomic weight appears to be too large, the atomic weights of the 
 elements naturally place them in this position. 
 
 Atomic Absolute Critical 
 Symbol Density weight critical pressure 
 
 temp. (mm.) 
 
 Helium (sun) He 2.0 4.0 
 
 Neon (the new one) Ne 10 . o 20.0 
 
 Argon (the lazy one) A 19.9 39.9 155.6 40,200 
 
 Krypton (the hidden one) .. Kr 40.9 81.8 210.5 41,240 
 
 Xenon (the stranger) Xe 64.0 128.0 287.8 43,500 
 
 " Critical data" from Travers, Experimental Study of Gases. 
 
 Ramsay gives the following table as representing an estimate of the 
 quantities in which the individual gases of the helium family are present 
 in the atmosphere. The data was obtained by distillation of " atmos- 
 phere" air, or argon. 
 
 Helium i to 2 parts per 1,000,000 of air. 
 
 Neon i to 2 parts per 100,000 of air. 
 
 Argon -937 part per 100 of air. 
 
 Krypton i part per 1,000,000 of air. 
 
 Xenon i part per 20,000,000 of air. 
 
 Expressed in parts per 1000: 
 
 Argon 9.37 
 
 Neon o.oi 
 
 Helium o . ooi 
 
 Krupton o . ooi 
 
 Xenon o . 00005 
 
 It is highly probable that the helium quickly passes out from our at- 
 mosphere into space, as a gas of so low a density must have, according 
 to the "kinetic theory" of gases, a velocity greater than the "critical 
 velocity" at the earth's surface which has a calculated value of 6.9 miles 
 per second. "The quantity of helium," says Travers, "which is con- 
 
198 EXPERIMENTAL CHEMISTRY. 
 
 stantly being given off by mineral springs is, however, enormous, so that 
 it is probable that the amount present in the atmosphere does not tend 
 to diminish." 
 
 The student is earnestly requested to examine the following books in 
 connection with the work of this chapter. 
 
 Experimental Study of Gases. Travers. Gas Analysis Hempel 
 Dennis. Liquefaction of Gases. Hardin. The Gases of the Atmos- 
 phere. Ramsay. 
 
 PROBLEMS. 
 
 1. Dumas, in determining the composition of air by passing it over 
 heated copper and measuring the residual nitrogen, tabulated the follow- 
 ing data: 
 
 Weight of tube and copper before experiment 120.00 grams. 
 
 Weight of tube and copper after experiment 121.15 grams. 
 
 Weight of globe exhausted 852 . oo grams. 
 
 Weight of globe and nitrogen 855 . 85 grams. 
 
 Calculate the percentage of nitrogen and oxygen by weight. By 
 volume. 
 
 2. Calculate the weight of 5 1. of air. Of 4.2 1. Of 9 1. 
 
 3. A mixture of 25 cm. 3 of air and 50 cm. 3 of hydrogen was exploded, 
 and the residue measured 60.3 cm. 3 What was the percentage of oxygen 
 in sample of air? 
 
 4. A student introduced 50.6 cm. 3 of air into a Hempel absorption 
 pipette containing alkaline pyrogallol. After agitating the pipette for 
 20 min. the gas was returned to the Hempel burette, and found to occupy 
 40.03 cm. 3 when measured under original conditions. Which component 
 of the atmosphere was removed ? Calculate the percentage of this com- 
 ponent by volume. 
 
 5. Dumas and Boussingault in 1841 found 12.373 grams of nitrogen 
 and 3.68 grams of oxygen in a sample of air. What per cent, of each 
 component did they find? 
 
 6. A U-tube containing phosphorus pentoxide was found to weigh 
 30.6293 grams. A volume of air which weighed 30.4268 grams was 
 passed through, after which the weight of the tube was found to be 
 31.0517 grams. Calculate the percentage of moisture present in the air. 
 
 7. It has been found that when air dissolves in water, the ratio of 
 oxygen to nitrogen in the dissolved air is no longer 1:3.71 (approx.) 
 tut i : ? The absorption coefficients of oxygen and nitrogen for water 
 as given by Bunsen are, O = .0411 and N = .0203, at 760 mm. Cal- 
 culate the ratio of free oxygen to nitrogen in the water when the total 
 atmospheric pressure is 760 mm. Hint. Recall Henry's Law and 
 Dalton's Law of Partial Pressures. 
 
 8. What properties manifested by the atmosphere led men to believe 
 that it was a chemical compound? 
 
 Enumerate the reasons which lead us to regard the air as a mixture. 
 
CHAPTER XIX. 
 
 NITROGEN FAMILY. 
 
 Nitrogen N, 14.01 
 
 Phosphorus P, 31.0 
 
 Arsenic As, 75.0 
 
 Antimony Sb, 120.2 
 
 Bismuth Bi, 208 . o 
 
 These five elements are grouped together not only because of their 
 atomic weight relationships, but in virtue of a similarity in chemical 
 properties. Antimony and bismuth, especially the latter, possess de- 
 cidedly metallic characteristics. With the exception of bismuth, they 
 form poisonous gaseous derivatives with three atoms of hydrogen. The 
 chief oxygen compounds of each are of R 2 O 3 and R 2 O 5 types. They form 
 generally unstable derivatives by combining with three atoms of a halogen. 
 These halogen compounds are particularly unstable in the presence of 
 water. It should be noted that with increase of atomic weight, the prop- 
 erties of the oxides of these elements vary gradually from strong acid 
 anhydrides to weak basic oxides. 
 
 NITROGEN, N. 
 
 At. Wt. 14.01 Mol. Wt. 28 02. 
 Preparation and Properties oj Nitrogen. 
 
 Experiment I. Recall Exps. I and II. "Nitrogen and the Atmo- 
 sphere." 
 
 Hydrogen derivatives, NH 3 , (NH 2 ) 2 , N 3 H. 
 
 Experiment II. Preparation and Properties of Ammonia Gas and Am- 
 monium Hydroxide. 
 
 (a) Preliminary Experiment. Dissolve about i gram of ammonium 
 chloride, NH 4 C1, in a few cm. 3 of water. Boil the solution. Do you 
 detect the odor of ammonia ? Test the escaping vapors with moistened 
 pieces of red litmus paper and turmeric paper. Results? Add a few 
 drops of a NaOH solution. Repeat previous tests. Results? Hold a 
 glass rod which has been moistened with strong HC1 just above the 
 mouth of the test tube. Results? What is the name of the substance 
 formed by the interaction of the ammonia gas and HC1? Equations? 
 Repeat, using KOH instead of NaOH. Other salts of ammonia may be 
 substituted for NH 4 C1. 
 
 109 
 
2OO EXPERIMENTAL CHEMISTRY. 
 
 (b) Laboratory method for preparation of ammonia gas. Support a 
 small flask by use of wire gauze on a ring-stand; fit a stopper provided 
 with a right-angled delivery tube to the flask; to the delivery tube con- 
 nect by means of rubber tubing a right-angled tube which is turned up- 
 ward that the gas may be collected by upward displacement of the air. 
 
 Mix together in a mortar about 20 grams of calcium oxide, CaO (quick- 
 lime), or calcium hydroxide, Ca(OH) 2 , and 10 grms of NH 4 C1. Odor? 
 Place the mixture in the flask and heat very gently; lay a piece of card- 
 board with a hole in it over the mouth of a dry bottle; invert the bottle 
 over the delivery tube. When the bottle is filled with gas, test by wav- 
 ing air from the bottle to the nose or by placing a piece of red litmus paper 
 moistened with water in the mouth of the bottle; cover it and place it 
 with mouth down upon the desk. Collect four bottles of the gas. Turn 
 the mouth of the delivery tube down, allowing it to barely touch the sur- 
 face of 10 cm. 3 of water in a test tube. Is there any evidence that the 
 gas is dissolving in the water ? Raise the test tube cautiously until the 
 mouth of the delivery tube is about 2 cm. below the surface of the water ? 
 Do bubbles of ammonia gas escape from the water? Explain. Do you 
 notice any change in the temperature of the water in the test tube ? 
 
 Lower the tube to its original position and continue to heat flask for 
 a few minutes, then remove the aqueous solution of the gas and extin- 
 guish the flame. Raise the flask from the gauze and place it on a piece 
 of dry cloth and allow it to cool. Explain the purpose of this last opera- 
 tion. When the flask cools, observe if there is any evidence of water in 
 the generating flask. Explain. Equations? 
 
 (c) Properties. In view of the method used in collecting the gas, w r hat 
 are your conclusions as to the density of ammonia gas as compared with 
 that of the air? Test the inflammability of the gas. Results? Will 
 the gas support combustion? (The gas w r ill burn in an atmosphere of 
 pure oxygen. Pass a stream of pure oxygen through concentrated solu- 
 tion of ammonia gas heated in a 50 cm. 3 flask. Apply a light to the mouth 
 of the flask. Explain. Equations?) Test the gas with moist turmeric 
 paper and with litmus paper. Results ? Moisten a glass rod with concen- 
 trated HNO 3 ; insert it into a bottle of the gas. Results? Equations? 
 
 (d) Solubility of the gas. Ammonium hydroxide. Place a bottle of 
 the gas mouth downward in a vessel of water. Hold it in this position for 
 4 or 5 minutes. Results. Explain. Examine the solution of ammonia pre- 
 pared in (b). Test it with litmus paper. Hold a rod dipped in concen- 
 trated HC1 over the solution. Results? What name is applied to this 
 aqueous solution of ammonia ? Indicate its formation by use of an equa- 
 tion. 
 
 (e) Stability of ammonium hydroxide. Put 5 cm. 3 of NH 4 OH in an 
 evaporating dish and boil it gently for a few minutes. Note the escape- 
 of the ammonia gas. Place 5 cm. 3 of NH 4 OH in a small beaker and 
 allow it to stand exposed to the atmosphere for 24 to 48 hours. Is 
 the odor as strong as before ? What do you infer as to the stability of 
 NH 4 OH? 
 
NITROGEN FAMILY. 2OI 
 
 (/) Does NH 4 OH manifest the properties of a base or an acid ? Con- 
 trast it with the aqueous solutions of hydrogen compounds of sulphur. 
 Are the basic properties of NH 4 OH due to the ammonia gas or the hy- 
 droxyl group formed by its combination with water ? Recall the electrical 
 conductivity of NH 4 OH. Is NH 4 OH a strong base? State the reasons 
 for your answer. What specific name do we apply to NH 4 ? Why? 
 Write the structural formula for NH 3 and NH 4 OH. 
 
 N + H 3 NH 3 + 1 1,900 cal. 
 NH 3 + Aq NH 3 ,Aq + 8500 cal. 
 
 Ice will melt in gaseous ammonia. Explain. 
 
 Experiment III. Production of Ammonia by Destructive Distilla- 
 tion of Nitrogenous Substances. 
 
 Heat separately in a test tube small quantities of each of the following: 
 hair, wool, bone, gelatin, feathers. Observe the effect of heat upon 
 these substances. Test the action of the fumes which are evolved upon 
 turmeric paper or red litmus paper. Results? Define destructive 
 distillation ? 
 
 Experiment IV. Preparation of Ammonium Salts. 
 
 (a) Preparation of ammonium chloride. Place 5 cm. 3 of NH 4 OH 
 (shelf-reagent) in a beaker; neutralize the base with dilute HC1, using 
 litmus paper or phenolphthalein as an indicator; evaporate the solution 
 until crystals begin to form on the sides of the vessel, then pour the solution 
 into a crystallizing dish and observe the formation of crystals as the 
 solution cools. Write the formula for the substances formed. 
 
 (b) Preparation of ammonium sulphate. Proceed as in (a), substituting 
 dilute H 2 SO 4 for HCL Are these salts soluble in water? Are they 
 conductors of electricity? If so, indicate by equations the nature of the 
 electrolytic dissociation. 
 
 Experiment V. Decomposition of Ammonium Salts. Dissociation. 
 
 (a) Place a small quantity of NH 4 C1 in a piece of hard glass tubing 
 open at both ends. Clamp the tube in a slightly inclined position; heat 
 the tube strongly and hold moist pieces of each kind of litmus paper in 
 each end of the tube. Results? Explain. Recall Exp. III. "Chemical 
 Equilibrium." 
 
 (b) Repeat (a), using (NH 4 ) 2 SO 4 instead of NH 4 C1. 
 
 Nitrogen and the Halogens. 
 
 Experiment VI. Preparation of an Endothermic Compound. Nitro- 
 gen Tri-iodide. 
 
 To a small quantity of powdered iodine in an evaporating dish, add a 
 little strong NH 4 OH. Frequent stirring will assist the digesting of the 
 iodine by the NH 4 OH. Let the mixture stand for half an hour. Filter. 
 
202 EXPERIMENTAL CHEMISTRY. 
 
 Place the black sediment upon several filter papers. Spread the wet 
 papers at a distance from one another and allow them to dry. The 
 dry black powder contains impure NI 3 , which is extremely explosive. 
 When touched with a feather it readily explodes. The shock produced 
 by the tread of a fly, or falling dust particles sometimes cause it to ex- 
 plode. The student should prepare only very small quantities of the sub- 
 stance. 
 
 (1) 2 NH 4 OH + I 2 -> NH 4 IO + NH 4 I + H 2 O. 
 
 (2) 3 NH 4 IO->N 2 H 3 I 3 + KOH. 
 
 (3) N 2 H 3 I 3 NH 3 + NI 3 . 
 
 (N,C1 3 ) = 38,100 cal. 
 
 Experiment VII. (Quant.) Determination of the Weight of a Liter 
 of Ammonia Gas. 
 
 Fill a 250 cm. 3 flask with the gas by upward displacement of air. Both 
 flask and gas must be perfectly dry; the gas may be dried by passing it 
 through two tubes filled respectively with small pieces of lime and soda- 
 lime. A loose plug of cotton should be placed in the neck of the flask 
 to prevent diffusion. When the flask is filled, slowly withdraw the tube 
 and cork at once. Wipe the flask and place it in the balance-room. Read 
 the thermometer and the barometer. Weigh the flask. Determine the 
 volume of the flask. Calculate the weight of this volume of air, if i cm. 3 
 of air at o C. and 760 mm. weighs .001293 gram subtract this weight from 
 the weight of the flask when filled with air. The remainder is the weight 
 of the vacuous flask. What is the weight of one liter of the gas at o C., 
 760 mm. ? Of the molar volume ? How many times heavier is one liter 
 of ammonia gas than a liter of hydrogen? What is the ratio of their 
 molecular weights ? 
 
 Experiment VIII. (Qual.). Determination of the Composition of 
 Ammonia Gas. 
 
 Instructions. Perfectly dry ammonia gas is passed very slowly 
 over heated magnesium powder placed in a piece of combustion tubing. 
 Loose plugs of cotton are placed in the end of the drying tubes contain- 
 ing, respectively, quicklime and soda-lime. The gaseous product is col- 
 lected over dilute H 2 SO 4 placed in a small glass dish. The greenish pow- 
 der which is formed in the combustion tubing is tested by pouring a por- 
 tion of it into a test tube half filled with water. Test the gas which is 
 liberated. Is it ammonia? Equation? 
 
 It is suggested that the ammonia gas be prepared by heating con- 
 centrated NH 4 OH over a low flame in a generating flask provided 
 with a thistle tube. The flask should rest upon a sand bath or a sheet of 
 asbestos. 
 
 Caution. Do not heat the thin layer of magnesium powder until the 
 apparatus is flooded with ammonia gas. 
 
NITROGEN FAMILY. 203 
 
 Experiment IX. (Quant.) L.T. Composition of Ammonia Gas. 
 (Volumetric.) 
 
 A very desirable form of the apparatus necessary for the performance 
 of this experiment may now be secured from any of the chemical labora- 
 tory supply houses. It is usually listed with the "Hofmann Lecture 
 Apparatus." A long glass tube, sealed at the lower end, and provided 
 at the upper end with a glass stop cock, which communicates above with 
 a small funnel-like chamber, which can be stoppered, is rilled with chlo- 
 rine over a saturated solution of sodium chloride. The tube is allowed to 
 stand mouth downward in the solution for sometime in order to let the 
 liquid drain out of it, after which the stop-cock is closed and the tube 
 on removal from the solution is placed in an upright position. The small 
 chamber above the stop cock is now nearly filled with concentrated ammo- 
 nium hydroxide and the stopper inserted. Slightly open the stop cock 
 and allow the ammonium hydroxide to pass into the lower tube drop by 
 drop. The evolution of considerable heat and usually a faint flash of 
 light, together with the formation of dense white fumes, indicate to the 
 experimenter that chemical reaction is taking place. Great care must be 
 exercised to prevent the escape of any gas or the entrance of any air into 
 the tube when the stop cock is opened. After nearly all of the ammonium 
 hydroxide has been allowed to pass into the tube, fill the funnel as before 
 and permit the liquid to pass gradually into the tube. Close the stop 
 cock. Fill a large beaker and the funnel w r ith dilute sulphuric acid; fit 
 a tube bent twice at right angles into a cork; fill the bent tube with the acid 
 solution and fit the cork to the funnel; place the other end of this bent tube 
 in the beaker of acid. The long tube should now be immersed in a tall 
 cylinder filled with water at the temperature of the laboratory. Open 
 the stop cock and let the acid drain slowly into the lower tube; if the 
 operation has been successful, the acid will flow into the tube until it is 
 two-thirds full. What is the residual gas ? Test it with a lighted match. 
 What are your conclusions in regard to the volumetric relations of the 
 two constituents of ammonia gas ? Write the equations for the reactions 
 involved in the above experiment. 
 
 NH 3 ,Aq + 3 Cl-* 3 HCl,Aq + N + 97,300 cal. 
 
 20,600 cal. 3 X 39, 300 cal. 
 
 Experiment X. (Quant.) L.T. Volumetric Composition of Ammonia 
 Gas. Chemical Equilibrium. 
 
 (a) Collect about 20 cm. 3 of pure dry ammonia gas in a perfectly 
 dry eudiometer by displacement of mercury. Allow the apparatus to 
 remain undisturbed for 10 or 15 min. The eudiometer should be clamped 
 in a vertical position in the trough of mercury (Fig. 36). Take the read- 
 ing of the thermometer suspended in the mercury, and record it together 
 with the barometric reading and the height of the mercury in the 
 eudiometer. Reduce the volume of gas to standard conditions. 
 
204 
 
 EXPERIMENTAL CHEMISTRY. 
 
 Connect the platinum wires of the eudiometer with an induction coil 
 and pass sparks through the gas for 30 min., or until the volume of the 
 gas is practically constant. Allow the apparatus to remain undisturbed 
 for 10 or 15 min. then tabulate data as before. Reduce this final volume 
 to standard conditions. State the ratio of the original and final volumes. 
 Explain, indicating the involved reactions by means of equations, (b) or (c) 
 may be performed at the discretion of the instructor. 
 
 (b) Continuation of (a). The relative volume of hydrogen and nitro- 
 gen may be determined by admitting sufficient dry oxygen to combine 
 with all of the hydrogen and exploding the mixture. The remaining 
 
 volume, of course, is nitrogen, ignoring aqueous tension. It is customary 
 to admit 5-10 cm. 3 of oxygen in excess of that required to combine with 
 the hydrogen. 
 
 (c) Chemical equilibrium. Continuation of (a). If a little sulphuric 
 acid is admitted above the mercury in the eudiometer and the mixture 
 of gases "sparked," the action will be reversed and the volume of the 
 mixture will gradually decrease until all of the gas disappears. Explain. 
 Interpret the following equations: 
 
 (2%-5%)2NH 3 <=N 2 
 2 NH 3 + N 2 
 
 (95% - 
 
 , -*(NH 4 ) 2 S0 4 . 
 
 Oxygen Derivatives. 
 
NITROGEN FAMILY. 205 
 
 Experiment XI. Preparation and Properties of the Oxides of Nitrogen.. 
 
 (a) Nitrous oxide (nitrogen monoxide), N 2 O. Place about 10 grams 
 of ammonium nitrate, NH 4 NO 3 , in the large test tube used for generating 
 oxygen; provide the tube with a cork which is fitted with a delivery tube. 
 Heat the test tube gently. Collect three bottles of the gas over warm 
 water by displacement of water. Observe the color and odor of the gas. 
 Is it inflammable ? Ascertain if it will support combustion by lowering a 
 glowing splint into a bottle of the gas. Results ? What other gas does 
 it resemble in its properties ? How can they be distinguished ? Explain 
 the action of nitrous oxide, N 2 O, in supporting combustion. Equations ? 
 
 (1) (N 2 ,0) = 17,700 cal. 
 
 (2) (N a O,O) = 24,800 cal. 
 
 In view of the above experiments and equation, would you conclude 
 that N 2 O is stable or unstable ? Is it an endothermic or exothermic com- 
 pound? Indicate the nature of the decomposition of NH 4 NO 3 by 
 means of an equation. 
 
 Note. Equation (2) reveals the reason why N 2 O is not reactive; i.e., 
 does does not combine with oxygen to form the higher oxides. 
 
 (b) Nitric oxide (nitrogen dioxide), N 2 O 2 and NO. Place about 15 
 grams of copper turnings in a generating flask provided with a thistle 
 tube and a delivery tube. Add 15 cm. 3 of dilute HNO 3 , then add 
 strong HNO 3 until there is a brisk evolution of gas. The addition of a 
 little cold water will diminish the speed of the reaction. When the brown- 
 ish fumes have been swept out of the apparatus, fill four bottles with the 
 gas by water displacement. What is the color of the gas ? Lift one of 
 the bottles from the water so as to admit air. What takes place ? What 
 is formed ? Is the gas inflammable ? Will it support combustion ? Test 
 by adding a few drops of CS 2 to a bottle filled with NO; shake well, then 
 bring a flame near the mouth of the vessel. Results ? Explain. Would 
 you infer that NO is stable or unstable ? Pass the gas into warm concen- 
 trated HNO 3 . Notice the formation of N 2 O 3 and N 2 O 4 which impart 
 a brown, green or blue color to the solution. Equation? Repeat, 
 using a dilute solution of potassium permanganate, KMnO 4 , acidified 
 with dilute H 2 SO 4 . Reoults? Equations? Does the gas act as an 
 oxidizing or a reducing agent? Pass a current of NO into a strong 
 solution of ferrous sulphate, FeSO 4 . Results? Equation? 
 
 (1) (N,O) = 21,500 cal. 
 
 (2) (N 2 ,0 4 )= 2,700 cal. 
 
 (3) ( 2 NO,0 2 )= 40,500 cal. 
 
 (4) (N0,0) - 13,450 cal. 
 
 (5) (2NO 2 ,O) = 2,800 cal. 
 
 After inspecting the above equations, what are your conclusions as 
 
2O6 EXPERIMENTAL CHEMISTRY. 
 
 regards the readiness with which NO combines with oxygen, and con- 
 versely, the ability of these higher oxides to oxidize ? 
 
 (c) Nitrogen trioxide, N 2 O 3 . 
 
 Place about one gram of arsenious acid (arsenious oxide), As 2 O 3 , in a 
 test tube; add 8 cm. 3 of strong nitric acid (1.30-1.35 sp. gr.) to cover 
 the solid, then heat gently. Note the color and odor of the evolved gas, 
 N 2 O 3 , as it is conducted into a test tube half filled with H 2 O Is the 
 gas soluble? Test the action of the solution on litmus paper. 
 
 Note. It is likely that a small quantity of nitrogen tetroxide, N 2 O 4 , is 
 produced simultaneously. Save the solution for use in Exp. XVI. 
 
 N 2 O 3 is the anhydride of what acid? See nitrous acid. Indicate by 
 equations the reasons for you answer. 
 
 (N 2 ,O 3 ) = 23,000 cal. 
 N 2 3 -N0 2 + NO. 
 
 (d) Nitrogen tetroxide (nitrogen peroxide), N 2 O 4 and NO 2 . 
 
 Heat about 8 grams of dry lead nitrate in a test tube. Note the color 
 and odor of the gas evolved. The vapor may be condensed in a U-tube 
 surrounded by a freezing mixture, or a portion of it may be conducted 
 into a concentrated solution of NaOH contained in a test tube. If all 
 of the gas is not absorbed by the solution, test the escaping gas for oxy- 
 gen. Results ? What is the name of the gas ? Examine the substance 
 which remains in the ignition tube. What is it? What is the name 
 and formula of the evolved gas? Write equations for all reactions. 
 Keep the NaOH solution for Exp. XVII. Label it. 
 
 (N,0 2 ) = -8, 1 25 cal. 
 
 (e) Nitrogen pentoxide (nitric anhydride), N 2 O 5 , does not exist in 
 the free state. It is usually prepared by distilling a mixture of phosphorus 
 pentoxide, P 2 O 5 , an d nitric acid. It forms colorless, rhombic prisms 
 which are so unstable that they explode violently when heated quickly. 
 They dissolve readily in water with the disengagement of considerable 
 heat, forming nitric acid, HNO 3 . 
 
 2 HN0 3 + P 2 5 N 2 5 + 2 HP0 3 . 
 (N 2 ,O 5 gas) = 12, ooo cal. 
 N 2 O 5 + H 2 O 2 HNO 3 . 
 
 The series of oxides of nitrogen are an excellent illustration of what law 
 of combination? State the law. Which of the oxides are colored? 
 Colorless ? 
 
 (H,N,O 3 liq.) = 41, 500 cal.; (H,N,O 3 Aq.) = 49,000 cal. 
 
 From the foregoing equations it is readily seen that all of the oxides of 
 nitrogen are endothermic compounds, i.e., the heat of formation from the 
 elements is negative. It is obvious then, that the nitrogen oxides cannot 
 
NITROGEN FAMILY. 207 
 
 be prepared from the elements without the addition of energy. (On 
 account of this energy relation, the atmosphere is able to preserve its 
 integrity as a mixture of nitrogen and oxygen gases, and this regardless 
 of the heat generated by frequent electrical discharges in the atmos- 
 phere.) These endothermic oxides are unstable, and even explosive. 
 Berthelot exploded nitric oxide by inflaming fulminating mercury. 
 
 An interesting bit of theory is associated with the thermal relations of 
 nitrogen and oxygen. It has been ascertained that the heat of combus- 
 tion of carbon or phosphorus in nitric oxide (NO) is about 21,500 calo- 
 ries larger than when the combustion takes place in oxygen. This is 
 usually explained upon the theory that the energy-content of nitric oxide 
 is greater than that of molecular oxygen (O 2 ). It is evident that the 
 molecules of the gas supporting the combustion must be split into their 
 constituent atoms. This operation will be attended by the absorption 
 or disengagement of heat. It follows then, that less heat is required to 
 separate NO into N and O than to separate O 2 into oxygen atoms. This 
 is interpreted as a proof that the molecules of free oxygen, as of other ele- 
 ments, possess an atomic structure. That is, the negative heat of for- 
 mation of nitric oxide (N,O = 21,500 cal.) indicates that the mutual 
 affinity of the nitrogen and oxygen atoms is less than the sum of the affin- 
 ities of the oxygen atoms for themselves and the nitrogen atoms for them- 
 selves in their respective molecules. 
 
 By inspecting these same equations it will be observed at once that al- 
 though the oxides of nitrogen are formed from the elements with an ab- 
 sorption of heat, the formation of the higher oxides from nitric oxide is 
 accompanied with a disengagement of heat. This accounts for the 
 readiness with which NO and N 2 O 3 tend to undergo oxidation, and con- 
 versely, the easy reduction of these higher oxides to nitric oxide accounts 
 for their marked oxidizing properties. (It will be recalled that these ox- 
 ides may act as oxidizing or reducing agents. See Exp. XL) On the 
 other hand, nitrous oxide (N 2 O) unites with oxygen to form the higher 
 oxide, nitric oxide (N 2 O,O), with an absorption of heat therefore its 
 slight reactivity with oxygen. Its oxidizing properties seem to be due 
 wholly to its instability. 
 
 In concluding this note, the student is reminded that the heat of for- 
 mation of compounds from the elements, as used in this book, is really 
 the heat of reaction (see note on Energetics of Chemistry), and may be 
 regarded not as an absolute but as a relative measure of the chemical 
 energy of the elements of which they are composed. To illustrate by 
 another example: Hydrogen and oxygen unite to form water as indi- 
 cated by the following equation: 
 
 2H 2 + O 2 2H 2 O gas + 2 X 57,061 cal. 
 
 Interpreting the equation in the light of our theories, a definite quantity 
 of heat will be required to split the molecules of each gas into its constit- 
 uent atoms; therefore, the quantity of heat which it is possible to measure 
 by calorimetric processes will be the true heat of union (heat of formation) 
 
208 
 
 EXPERIMENTAL CHEMISTRY. 
 
 less this amount. This observed heat, which is really the heat of reaction, 
 merely indicates that the mutual affinity of hydrogen and oxygen is greater 
 than the sum of the affinities of the hydrogen atoms for themselves and 
 the oxygen atoms for themselves in their respective molecules, by an 
 amount, the thermal equivalent of which is equal to 57,061 calories. 
 All affinity calculations derived from thermal data by such processes 
 must be purely relative. 
 
 Note. The heat of dissociation of the hydrogen molecule into atoms, 
 H 2 (H,H), has a recorded value of 128,000 cal. (Richter.) 
 
 Experiment XII. Preparation of Nitric Acid. 
 
 Place ten grams of pulverized sodium nitrate, NaNO 3 in a small tubu- 
 lated retort (Fig. 37) or a distilling flask; support the retort on a wire gauze 
 in such a manner that it can be heated conveniently. Introduce the end of 
 
 FIG. 37. 
 
 the retort into a receiver (a dry test tube will do) which is immersed in cold 
 water. Add 20 cm. 3 of strong H 2 SO 4 by aid of a funnel that no acid may 
 enter the neck of the retort. Insert the stopper in the retort. Is there 
 any evidence of chemical action taking place? Heat the retort gently 
 and observe the gradual accumulation of liquid in the test tube. What 
 is its color and odor? Action on litmus paper? Allow a few drops of 
 the liquid product to run down the inside of a test tube containing 
 a solution of ferrous sulphate, FeSO 4 . Results? Equations? Repeat 
 above tests using the shelf-reagent, HNO 3 . Results? What is your 
 
NITROGEN FAMILY. 2OQ 
 
 conclusion as to the identity of the liquid in the receiver? What is the 
 substance in the retort? Write the equation for the interaction of Na- 
 NO 3 and H 2 SO 4 . Which is the more volatile, H 2 SO 4 or HNO 3 ? 
 
 Assuming that the reaction is reversible, what is the effect of heat 
 upon the equilibrium ? 
 
 Experiment XIII. Properties of Nitric Acid. 
 
 (a) Action of nitric acid on bases. To a beaker of water add a few 
 drops of NaOH or KOH; test with red litmus paper. Now add a dilute 
 solution of HNO 3 , drop by drop, until neither blue nor red litmus paper 
 is altered in color when dipped into the solution. Taste a drop of the 
 solution. What has been formed by the interaction of NaOH and HNO 3 ? 
 Write the "ionic" equation for the interaction. 
 
 (6) Oxidizing action. Test its action on indigo. Equation ? Recall 
 the action of concentrated HNO 3 on iodine. Equation ? Try the action 
 of strong boiling HNO 3 on powdered sulphur. Is there any evidence of 
 chemical action ? Dilute the solution, and test for a sulphate. Results? 
 Equations? To a solution of ferrous chloride or sulphate add a little 
 NH 4 OH. Result? Allow contents of tube to remain exposed to the air. 
 Results? Repeat using ferric chloride. Results? To a solution of 
 ferrous sulphate add a few drops of concentrated HNO 3 ; heat gently and 
 then add NH 4 OH. Results? Explain the action of HNO 3 upon the 
 " ous " salt. Equations ? What property of HNO 3 which is particularly 
 characteristic of it, is emphasized in these reactions ? 
 
 (c) Reduction of HNO 3 . Formation of ammonia. Place a few grams 
 of zinc dust in a test tube, and add a solution of KOH (i of KOH to 20 
 of H 2 O). Warm the mixture gently, and test the evolved gas. What is 
 it ? Add a few drops of dilute HNO 3 . Note the odor of the gas which is 
 now given off. Confirm your conclusions by using two other tests. 
 Results? Explain. Equations? 
 
 (d) Try the action of both dilute and concentrated HNO 3 upon each 
 of the following metals: zinc, iron, copper, lead and tin. Note in each 
 case the kind of gas evolved. Be on the alert for hydrogen, oxides 
 of nitrogen and ammonia. Devise a method for showing whether the 
 product of the reaction remaining in the test tube in each case is a nitrate 
 or not. Apply it. Equations? Compare the interaction of HNO 3 and 
 the metals with its action toward the non-metals. Refer to lecture notes 
 and reference texts for assistance in balancing equations. 
 
 (e) Action on nitrogenous substances. Dip a quill, a piece of wool 
 or pieces of white feather, in concentrated nitric acid. The yellow color 
 is due to xanthoproteic acid, formed by the interaction. This is regarded 
 as a test for nitrogen. 
 
 (/) Determine the relative electrical conductivity of a 5N solution of 
 HNO 3 . Results? What is its percentage of dissociation in a iN 
 solution? (Refer to tables.) Would you conclude that the acid is very 
 active or relatively inactive? Why? 
 
 Enumerate several chemical properties of HNO 3 which are especially 
 14 
 
210 EXPERIMENTAL CHEMISTRY. 
 
 characteristic. Structural formula for nitric acid? Anhydride of the 
 acid? 
 
 Note. The complexities of oxidation by nitric acid are of such nature 
 that the student usually requires some assistance in mastering the essential 
 principles of the reactions involved. The principles indicated in the 
 following equations should be thoroughly understood. The gaseous 
 products resulting from the interaction of HNO 3 and the metals and non- 
 metals are determined largely by the concentration of the acid. 
 
 For dilute nitric acid: 
 
 2 HNO 3 2 NO + H 9 O + (3 O) 
 S + (3 O) + H 2 H 2 S0 4 . 
 
 2 HNO 3 + S -> H 2 SO 4 + 2 NO. 
 
 For concentrated nitric acid : 
 
 3(2HNO 3 ) 2 NO 2 +' H 2 O + (O) 
 S + 3(0) + H 2 H 2 S0 4 
 
 6HNO 3 + S H 2 SO 4 + 2H 2 O + 6NO 2 . 
 
 Very dilute nitric acid: 
 
 Zn + 2 HNO 3 Zn(NO 3 ) 2 + (2H) 
 HNO 3 + 68(H) NH 3 + 3 H 2 O 
 or NH 3 + HN0 3 NH 4 NO 3 . 
 
 Dilute nitric acid: 
 
 2 HN0 3 ->H 2 + 2 NO + ( 3 0) 
 3 Cu + ( 3 0)-> 3 CuO 
 3CuO + 6HNO 3 3 Cu(NO 3 ) 2 + 3H 2 O 
 
 3 Cu + 8HNO 3 - Cu(NO 3 ) 2 + 4H 2 O + 2NO. 
 
 Concentrated nitric acid: 
 
 2 HNO 3 H 2 O + 2 NO 2 + (O) 
 Cu + (O) CuO 
 CuO -t- 2 HNO 3 -> Cu(NO 3 ) 2 + H 2 O 
 
 Cu + 4 HNO 3 -> Cu(NO 3 ) 2 + 2 H 2 O + 2NO 2 . 
 
 The oxidation of "ous" salts to "ic" salts by means of nitric acid may 
 be represented by equations as follows: 
 
 3 Fe(N0 3 ) 2 + 4 HN0 3 3 Fe(N0 3 ) 3 + 2 H 2 O + NO, 
 or 2 Fe(NO 3 ) 2 + 4 HNO 3 2 Fe(NO 3 ) 3 + 2H 2 O + 2 NO 2 . 
 
NITROGEN FAMILY. 211 
 
 According to the "ion theory:" 
 
 *HNO 3 <=NO' 2 + OH' 
 
 HNO 3 + H 2 O <= NO' ' ' + 3OH' 
 
 3 Fe + 6NO' 3 + NjD^+^OH' + 3 H" + 3 NO' 3 = 3 Fe 
 
 9 NO' 3 + 3 H 2 + NO. 
 
 or 2Fe ' ' + 4NO' 3 + 2NO 2 + OH' -f 2H' + 2NO' 3 = 2Fe + 
 
 6NO' 3 + 2H 2 O ^~ 
 
 It is obvious that oxidation is a process whereby the number oj positive 
 charges are increased, or the negative are decreased. Reduction consists oj 
 increase oj negative charges and a decrease oj the positive. 
 
 Experiment XIV. Preparation and Properties of Aqua Regia. Oxida- 
 tion of Hydrochloric Acid by Nitric Acid. 
 
 To one volume of concentrated HNO 3 add three volumes of strong HC1. 
 Gently warm the mixture and notice its appearance and odor. This 
 mixture 
 
 HNO 3 + 3HC1 2 H 2 O + NOC1 + Cl 2 , 
 
 of acids is usually referred to as aqua regia (royal water), because it will dis- 
 solve the noble metals, gold, platinum, etc., forming chlorides. No single 
 acid with the exception of selenic acid, will dissolve gold. Place a small 
 piece of gold leaf in each of two test tubes. To one add 2 cm. 3 of HNO 3 , 
 and to the other 6 cm. 3 of HC1. Observe that the gold leaf is not attacked. 
 Mix the contents of both test tubes and warm gently. Results? What 
 is formed ? Equations ? What is the relative solution tension of gold ? 
 (refer to table of solution-tensions). Explain the solubility of the " noble " 
 metals in aqua regia. 
 
 Experiment XV. Effect of Heat on Nitrates. Oxidizing Power of 
 Nitrates. 
 
 (a) Recall Exp. XI (a). Equation ? The nature of the decomposition 
 is peculiar to ammonium nitrate. 
 
 (b) Recall Exp. XI (d). Equation? The result is typical of the nitrates 
 of the heavy metals. 
 
 (c) Heat sodium nitrate in a hard glass test tube before the flame of a 
 blast-lamp if necessary to secure an evolution of gas. Test the gas for 
 oxygen. Continue to heat until gas is no longer evolved. What is 
 the composition of the residue ? (Do not attempt to answer this question 
 until you have performed Exp. XVI.) The action of sodium nitrate is 
 characteristic of the nitrates of the alkali metals and the alkali-earth 
 metals. Keep the residue for Exp. XVI. 
 
 Note. A residue of the same composition may be prepared more 
 easily by melting the nitrate with a piece of lead in a crucible, and keeping 
 * Ostwald, Grundriss der allgemeinen Chemie. 
 
212 EXPERIMENTAL CHEMISTRY. 
 
 the mixture stirred with an iron spatula. What is action of the nitrate 
 on the lead? 
 
 Experiment XVI. Preparation and Properties of Nitrous Acid. 
 
 (a) Nitrous acid, HNO 2 , is formed by the interaction of a nitrite, say 
 sodium nitrite, NaNO 3 , and H 2 SO 4 . 
 
 Add a small quantity of dilute H 2 SO 4 to a dilute solution of NaNO 2 . 
 Test for HNO 2 as follows: Add a few drops of the mixture supposed to 
 contain HNO 2 to a solution of KI, then add 2 cm. 3 of CS 2 and shake 
 vigorously. Free iodine proves presence of HNO 2 . Equations? 
 Prove that a nitrate will not give the same results. Record your data. 
 
 Try the action of the HNO 2 on a dilute solution of KMnO 4 . Results? 
 Equation ? Does HNO 2 act as an oxidizer or as a reducing agent relative 
 toKI? To KMnO 4 ? 
 
 (b) Test the aqueous solution of N 2 O 3 prepared in Exp. XI (c) for the 
 presence of HNO 2 . Results? Equations? What is the anhydride of 
 HNO 2 ? 
 
 (c) Dissolve the residue from heating NaNO 3 (Exp. XV) in a small 
 quantity of water. Filter. Add dilute H 2 SO 4 to the nitrate and test for 
 the presence of HNO 2 . Results ? Equation ? Was the residue a nitrate 
 or a nitrite? What is the effect of heat on NaNO 3 ? Record your 
 answer under (c), Exp. XV. 
 
 What class of salts are formed by the interaction of nitrous acid and 
 the bases? Write the structural formula for HNO 2 . 
 
 (H,N,O 2 Aq.) - - 30,700 cal. 
 
 Experiment XVII. Tests for Nitrates and Nitrites. 
 
 (a) Test for nitrates. Dissolve a small crystal of ferrous sulphate, 
 FeSO 4 , in 3 or 4 cm. 3 of water in a test tube. Add a small quantity of 
 the solution of any nitrate. Now hold the tube in a slanting position and 
 carefully pour down the side 3 or 4 cm. 3 of strong H 2 SO 4 . This should 
 be done in such a manner that the acid may form a layer at the bottom. 
 Notice the brown ring which forms at the boundary between the two 
 liquids. Explain. Equation ? 
 
 (b) Tests for nitrites. Repeat (#), substituting a nitrite for the nitrate. 
 Results? Can this test be used to detect a nitrite in the presence of a 
 nitrate, or vice versa? Recall the tests for HNO 2 (nitrite) used in Exp. 
 XVI. These tests will enable you to detect a nitrite in the presence of a 
 nitrate, but the nitrite must be removed before (a) may be applied as a 
 test for the nitrate. This is usually accomplished by boiling the solution 
 with NH 4 C1. See Exp. I for the reaction. 
 
 The tests are usually carried out as follows: 3 or 4 cm. 3 of the "un- 
 known" solution are placed in a test tube, acidulated with dilute H 2 SO 4 , 
 and tested for a nitrite, using the KI or the KMnO 4 test. If a nitrite is 
 absent, the solution is tested at once for a nitrate; but if a nitrite is present, 
 a portion of the "unknown" solution is boiled with NH 4 C1 until small 
 
NITROGEN FAMILY. 213 
 
 portions of the solution which are removed, give no test for the nitrite. 
 The test for nitrates is then applied. 
 
 (c) Prepare a solution containing a nitrite and a nitrate and proceed 
 with the analysis as per directions. Record all data. 
 
 (d) Test the NaOH solution of N 2 O 4 prepared in Exp. XI (d), for 
 nitrates and nitrites. Results? What do you infer was the action of 
 the N 2 O 4 on the NaOH solution? 
 
 All nitrates and nitrites are soluble in water. The silver salt of nitrous 
 acid is sufficiently difficult of solution to be precipitated on the addition of 
 silver nitrate to a solution of a nitrite if the latter is not too dilute. 
 
 PHOSPHORUS, P. 
 
 At. Wt. 31 Mol. Wt. (P X -P 4 ). 
 
 Phosphorus, like sulphur, exists in several allotropic forms crystalline 
 or yellow phosphorus, and amorphous or red phosphorus. The properties 
 of these forms differ widely although they are composed of the same 
 materials. However, their respective energy-contents are very different. 
 See "Note on the Energetics of Chemistry." 
 
 (Yellow) P > (Red) P = 27,300 cal. 
 
 The approximate thermal equivalent of the difference between the in- 
 trinsic energies of the allotropic forms is 27,300 cal. 
 
 Experiment I. Properties of Phosphorus. 
 
 Caution. Yellow phosphorus must be kept and cut under water. It 
 must be handled with forceps, never with the fingers. 
 
 Determine the physical properties of the yellow and red varieties of 
 phosphorus. Place a quantity of powdered boneblack upon an iron 
 plate (tile or brick). Lay a piece of yellow phosphorus about the size 
 of a pea upon this and heap the powder up around the phosphorus, but 
 leave the top of the phosphorus exposed to the air. Does the phosphorus 
 take fire? Explain. Equations? Repeat the experiment using red 
 phosphorus. Results ? Explain. 
 
 Determine the relative solubility of the two forms of phosphorus in 
 CS 2 . Results? Dissolve a very small piece of yellow phosphorus in 
 about i cm. 3 of CS 2 . Place a filter paper on the ring of the ring-stand; 
 pour the solution upon the filter paper. Let the CS 2 evaporate without 
 heating. Result ? Equation ? 
 
 Prove that yellow phosphorus is practically insoluble in water or alcohol. 
 
 Drop a small piece of yellow phosphorus into a flask half filled with 
 water. Heat the flask until the phosphorus is melted, then pass a current 
 of oxygen through a delivery tube upon the melted phosphorus in. the 
 bottom of the flask. Results ? 
 
 Put a drop of liquid bromine in a dry bottle; allow the bromine to 
 vaporize. When the bottle is filled with the vapor, throw a small dry piece 
 
214 EXPERIMENTAL CHEMISTRY. 
 
 of yellow phosphorus into the bottle (?). Equation? Is the product 
 stable? Repeat, using red phosphorus (?). 
 
 Bring together in a porcelain crucible or an evaporating dish a flake 
 of iodine and a very small dry piece of yellow phosphorus (?). What 
 is the source of the light and heat ? Equation ? Is the product stable ? 
 Does red phosphorus yield similar results ? 
 
 Place a small quantity of red phosphorus in the bottom of a test tube. 
 Clamp the tube in a horizontal position, and gently heat the end containing 
 the phosphorus. What is the yellow substance which collects on the 
 cold portion of the tube ? Verify your conclusion by suitable tests. Did 
 the red phosphorus melt or sublime? 
 
 Tabulate in vertical columns against one another the corresponding 
 properties of these two modifications of phosphorus. To what is this 
 difference of behavior attributed? 
 
 Hydrogen Derivatives. 
 
 Experiment II. (Hood) Preparation of Phosphine. 
 
 (a) Into a beaker nearly full of water (a few drops of HC1 will increase 
 the speed of the reaction) drop a small piece of calcium phosphide, Ca 3 P 2 
 (?). Equation? Recall the action of magnesium nitride, Mg 3 N 2 , 
 upon water. 
 
 Note. Calcium phosphide has an irregular composition, and because 
 of this fact, a mixture of the three hydrides, PH 3 , P 2 H 4 , P 4 H 2 is obtained. 
 Which of the three inflames spontaneously in the air and gives to gaseous 
 phosphine its spontaneous inflammability? By what other names is 
 gaseous phosphine known ? 
 
 (b) A small generating flask (100-250 cm. 3 ) is fitted with a rubber 
 stopper provided with two right-angled delively tubes, a and b (Fig. 38). 
 
 A short piece of rubber tubing, carrying 
 a pinch cock is attached to a. In the 
 flask, place a strong solution of KOH 
 (i of KOH to 2 of H 2 O). The flask 
 should be half filled, then place it on a 
 piece of iron gauze supported on the ring 
 stand in such manner that the flask may 
 be heated. Connect b with a delivery 
 - tube bent so that its lower end, which is 
 FIG. 38. (Smith and Keller.) turned upward, dips beneath the surface 
 
 of water in a large beaker or a pneumatic 
 
 trough. The water should have a temperature of about 20. Now drop 
 a piece of phosphorus about the size of a pea into the flask and stopper, 
 air-tight. Connect a with a hydrogen generator and pass a current of 
 hydrogen through the apparatus until all of the air is displaced; disconnect 
 from the generator and close pinch cock on a. (A current of coal gas may 
 be used instead of the hydrogen, or a few drops of ether may be added to 
 
NITROGEN FAMILY. 215 
 
 the contents of the flask. The ether on evaporating will drive the air out 
 of the flask.) Heat the flask gently and phosphine will be evolved (?). 
 Equation? What is formed by the burning of the phosphine? Equation? 
 
 Caution. Do not remove the lamp or the tube from the water until 
 the reaction has terminated, then lift tube from water first, and remove 
 the lamp. The gas is poisonous. 
 
 Read "Note" in (a) if you have not already done so. Is there any 
 similarity between ammonia and phosphine? 
 
 (Yellow) P + 3H PH 3 + 11,600 cal. 
 
 Phosphorus and the Halogens. We have seen that the halides of phos- 
 phorus may be formed by the direct union of the elements. The halides, 
 being exothermic compounds, are very much more stable than the analo- 
 gous compounds of nitrogen. In contact with water they are unstable, 
 undergoing hydrolysis, with the formation of a hydrogen halide and an 
 oxacid of phosphorus. The reactivity of the halides of phosphorus with 
 water is fully explained by the large amount of heat liberated at the time 
 of the action. 
 
 (P,C1 3 ) = 75,300 cal. (P,C1 5 ) = 104,900 cal. 
 
 (P,Br 3 ) = 42,600 cal. (P, I 3 ) = 10,900 cal. 
 
 Experiment III. Decomposition of the Halides of Phosphorus by 
 Water. Hydrolysis. 
 
 (a) Preparation of phosphorous acid by hydrolysis of phosphorus 
 trichloride. Place a few drops of phosphorus trichloride, PC1 3 , in a test 
 tube; blow your breath across the mouth of the test tube (?). Recall 
 the test for a hydrogen halide. Now add a small quantity of water to the 
 oily liquid ( ?). Ascertain whether any thermal phenomena accompanies 
 the dissolving of the PC1 3 or not ( ?). Warm the tube, and again blow your 
 breath across its mouth (?). What are your inferences? Equations? 
 Evaporate the solution to dryness on a water-bath until all of the HC1 
 has passed off. What is the residue ? Add a little water to the residue 
 and try the effect of the solution upon a AgNO 3 solution. A black pre- 
 cipitate proves the presence of phosphorous acid, H 3 PO 3 . 
 
 (b) Preparation of phosphoric acid by hydrolysis of phosphorus penta- 
 chloride. 
 
 By means of a spatula, place a small quantity of PC1 5 in a test tube. 
 Blow your breath across the mouth of the tube (?). Add a few cm. 3 of 
 water. Note the hissing noise which accompanies the reaction. Is the 
 test tube warmed by the interaction of the substances ? Warm the tube 
 and again blow your breath across its mouth (?). Test it with litmus 
 paper ( ?). Boil until the solution is free of HC1. What is the remaining 
 fluid? Add a few drops of it to a AgNO 3 solution (?). A yellow pre- 
 cipitate of silver orthophosphate, Ag 3 PO 4 , proves the presence of phos- 
 phoric acid, H 3 PO 4 . Equations ? 
 
2l6 EXPERIMENTAL CHEMISTRY. 
 
 Recall the methods used for the preparation of pure hydriodic and 
 hydrobromic acids. 
 
 Oxygen Derivatives. 
 
 Experiment IV. Preparation and Properties of Phosphorus Pentoiide. 
 Phosphoric Acid. 
 
 Burn a carefully dried piece of phosphorus under a bell- jar. Allow 
 the jar to remain undisturbed until the heavy white vapors have deposited 
 upon its sides. Compare the white powder with the substance labeled 
 "phosphorus pentoxide," P 2 O 5 . Place a portion of each in separate 
 test tubes containing a little water (?). Is there any thermal evidence of 
 chemical action? Test each with litmus paper (?). Compare their 
 actions toward a silver nitrate solution (?). When phosphorus burns 
 with a free supply of oxygen, what compound is formed? Formula? 
 When this compound reacts with water, what is the product ? Equation ? 
 Name two other oxides of phosphorus and give the formula for each. 
 What is the anhydride of phosphoric acid ? Of phosphorous acid ? 
 
 (P 2 A) = 369,800 cal. (P 2 5 , Aq.) = 41,600 cal. 
 
 Phosphorus pentoxide is one of the most effective drying agents known. 
 It has been shown that a glass tube, four or five inches in length, filled 
 with the phosphorus pentoxide, will entirely dry a gas which is slowly 
 passing through. 
 
 Experiment V. Preparation and Properties of Phosphoric Acid. 
 
 Phosphoric acid may be prepared by heating a very small quantity of 
 red phosphorus with an excess of strong nitric acid; filter and remove 
 excess of water by evaporation on the steam bath. The thick syrup 
 which remains should be dissolved and tested for phosphoric acid w r ith a 
 silver nitrate solution ( ?). 
 
 What is the formula for phosphoric acid? How many replaceable 
 hydrogen atoms ? Is it unibasic, dibasic or tribasic ? Into what ions does 
 it dissociate when placed in water. Is it a "strong" acid? (Refer to 
 tables.) State the reasons for your answer. Write the structural formula 
 for phosphoric acid. Give the names and formulae of three other acids 
 of phosphorus. 
 
 From data given in Exp. IV calculate the "heat of formation" of phos- 
 phoric acid in aqueous solution from its elements. 
 
 Experiment VI. Reactions of the Orthophosphates. 
 
 (a) Dissolve a little disodium hydrogen orthophosphate, Na^RPC^, 
 in water. Test a portion of the solution with litmus paper ( ?). Explain. 
 
 (b) To this portion add silver nitrate in solution ( ?). Filter. Ascer- 
 tain the solubility of the yellow precipitate in HNO 3 ( ?). In NH 4 OH ( ?). 
 Equations ? 
 
NITROGEN FAMILY. 217 
 
 (c) Test a portion of the original solution with ferric chloride, FeCl 3 ( ?). 
 Add sufficient HC1 to dissolve the precipitate, then add slowly NH 4 OH 
 until the solution is alkaline ( ?). Equations? 
 
 (d) Test a solution of any soluble orthophosphate with " magnesia 
 mixture" (?). Is the precipitate crystalline?* Is the precipitate soluble 
 in excess of MgSO 4 ? Is^it reprecipitated in the cold by NH 4 OH ? Equa- 
 tions ? 
 
 Note. "Magnesia mixture" is prepared by adding a little NH 4 OH 
 and an excess of NH 4 C1 to a solution of magnesium sulphate, MgSO 4 . 
 
 (e) Add an ammonium molybdate, (NH 4 ) 2 MoO 4 , solution to a dilute 
 solution of an orthophosphate acidified with concentrated HNO 3 ( ?). 
 Describe the precipitate. What is it ? Give the formula. Equations ? 
 
 Note. (d) and (e) are reactions frequently employed in analytical 
 chemistry. 
 
 Experiment VII. Preparation and Reactions of the Pyrophosphates. 
 
 When disodium hydrogen orthophosphate, Na 2 HPO 4 , is heated to 
 about 300, each two equivalents of it evolve one equivalent of water and 
 a neutral sodium pyro phosphate is formed as a colorless glassy mass, soluble 
 in water. 
 
 Place a small quantity of Na 2 HPO 4 in a crucible and heat to redness ( ?). 
 Equations ? Dissolve the glassy residue ( ?) in water and test the solution 
 as per Exp. VI., (a), (&), (c), (d), (e). Tabulate results. Equations? 
 
 Test a solution of H 4 P 2 O 7 (or a pyrophosphate acidified with acetic acid) 
 with a clear solution of albumen. Is the latter coagulated ? 
 
 Experiment VIII. Preparation and Reactions of the Metaphosphates. 
 
 (a). Heat to redness a little microcosmic salt (sodium ammonium 
 hydrogen orthophosphate, NaNH 4 HPO 4 ), and dissolve the residue of 
 sodium metaphosphate in water. Test the solution as per Exp.. VI. , 
 a, b, c, d, e. Tabulate results. Equations? 
 
 Test a solution of HPO 3 with a clear solution of albumen. Is the 
 latter coagulated ? 
 
 (b) Bend the end of a platinum wire into a closed loop; heat it to red- 
 ness, then plunge it into a little microcosmic salt; heat strongly in the outer 
 zone of the blow-pipe flame until a "glassy bead" forms ( ?). Touch the 
 hot bead to some powdered CuO and heat strongly again. If the opera- 
 tion has been successful a clear blue-green bead will be formed. The 
 solubility of certain metallic oxides in the metaphosphates and the meta- 
 borates is taken advantage of to detect their presence. The color of the 
 bead is characteristic of the oxide. 
 
 *If the precipitate formed above is not crystalline add a slight excess of HC1 and 
 reprecipitate with very dilute NH 4 OH. 
 
2l8 EXPERIMENTAL CHEMISTRY. ' 
 
 Experiment IX. Tests to Distinguish the Phosphates. 
 
 Mention one reaction which will enable you to distinguish orthophos- 
 phates from all other phosphates. Try it. Record method and results. 
 How would you distinguish between a meta- and a pyrophosphate ? 
 Try it. Record all data. 
 
 With the exception of members of the potassium family, the normal 
 orthophosphates and pyrophosphates of all the metals are insoluble in 
 water. 
 
 ARSENIC, AS. 
 
 At. Wt. 75.0 Mol. Wt. (As 2 , 6 4 5-As 4 , 1,700). 
 
 Arsenic is essentially an acid-forming element and is, therefore, a non- 
 metal. In some of its compounds, however, it functions apparently 
 as a metal. It cannot displace hydrogen from dilute acids. It has 
 valences of three and five. 
 
 Experiment I. Properties of Arsenic. 
 
 (a) Study the physical properties of the elementary substance. Does 
 it possess a metallic appearance ? Is it crystalline in form ? Place a 
 piece of arsenic about half the size of a grain of wheat in a small tube of 
 hard glass and heat to redness ( ?). What name is applied to such change 
 of state ? Repeat using orpiment, As 2 S 3 , or realgar, As 2 S 2 . 
 
 (b) Heat a very small piece of arsenic on charcoal in the oxidizing flame. 
 Results ? Note that the characteristic odor of garlic is perceptible. Will 
 the free element burn in the air? If so, what is formed? 
 
 (c) To a small quantity of arsenic in a test tube add an excess of HNO 3 . 
 Boil so long as brown fumes form, or until the liquid does not color on cool- 
 ing. Save the solution for use in Exp. V. 
 
 Recall similar experiments with iodine and phosphorus (?). 
 
 Hydrogen Derivatives. 
 
 Experiment II. Arsine. Marsh's Test for Arsenic. 
 
 (Hood; Poison! Instructions). Thoroughly clean the hydrogen- 
 generator; place a small quantity of chemically pure zinc in the flask; 
 connect the generator with a tube (U-tube or bulb-tube) filled with cal- 
 cium chloride, CaCl 2 . Connect this latter tube with a piece of hard glass 
 tubing, r, about 20 cm. long, clamped in a horizontal position (Fig. 39). It 
 is well to make constrictions in the tube at intervals of 7 or 8 cm. by gently 
 heating and drawing it out. Now add dilute HC1 through the thistle tube. 
 When the air has been displaced by hydrogen, light the gas by the "test tube" 
 method. Observe the color of the flame. By means of a pair of pinchers 
 hold a porcelain crucible lid in the flame. If there is no deposit of solid 
 matter upon the lid, introduce through the thistle tube a few drops of a 
 solution of arsenic trichloride (arsenic trioxide, As 2 O 3 , dissolved in HC1). 
 
NITROGEN FAMILY. 
 
 219 
 
 Note the appearance of the flame ( ?). Is there a deposit upon the cold 
 porcelain lid? Marsh's test: Heat the horizontal tube between any 
 two constrictions, with a flame. Results ? Locate the deposit in the tube 
 relative to generator and the flame. Has the deposit upon the lid a me- 
 tallic luster? Test the solubility of the deposit (arsenic "spot") with 
 any soluble hypochlorite. Indicate by equations the chemical reactions 
 
 FIG. 39. 
 
 involved in the preparation of the gas, arsine. Explain by use of equa- 
 tions the changes occurring within the flame. What other name is some- 
 times applied to the gas, AsH 3 ? 
 
 Note. The student may be asked to determine whether or not arsenic 
 is present in the green coloring matter used on wallpaper and shipping- 
 labels. 
 
 As + H 3 * AsH 3 iijoocal. 
 
 Oxygen Derivatives. 
 
 Experiment III. Arsenic Trioxide (Arsenious Oxide). Arsenious Acid. 
 
 (a) Place a little powdered arsenic in a hard glass tube open at both 
 ends and about 20 cm. long. Clamp the tube in a nearly horizontal posi- 
 tion and heat. Notice the white deposit of arsenic trioxide, As^g, on 
 the colder portions of the tube. Scrape the deposit into a small test tube 
 and resublime it. Does the sublimate show a crystalline structure? 
 By what other name is arsenic trioxide commonly known ? 
 
 (As 2 , O 3 ) = 1 54,600 cal. 
 (A&j, O 5 ) = 2 19,400 cal. 
 
 (b) Reduction of arsenic trioxide. Heat a pinch of As 2 O 3 after mixing 
 it with a little powdered wood charcoal, in a very narrow test tube (or, 
 
220 EXPERIMENTAL CHEMISTRY. 
 
 better, in a drawn-out glass tube having a small bulb on the end). Re- 
 sults ? Equation ? 
 
 Note. It is advisable to cover the mixture with a layer of powdered 
 charcoal. 
 
 (c) Preparation of arsenic trichloride. Boil a small quantity of As 2 - 
 O 3 with concentrated HC1 (?). Equation? Add a few cm. 3 of water. 
 Preserve the solution for future use. Does As 2 O 3 in this reaction manifest 
 the properties of a metallic or a non-metallic oxide, i.e., basic or acidic prop- 
 erties ? 
 
 (As, C1 3 ) = 71,400 cal. 
 
 (d) Preparation of sodium arsenite, Na 3 AsO 3 . Boil a small quantity 
 of As 2 O 3 with a solution of sodium hydroxide. Results ? Equation ? 
 
 Does As 2 O 3 in this reaction manifest the properties of a metallic or a 
 non-metallic oxide? 
 
 (e) Formation of arsenious acid, H 3 AsO 3 . Heat a little As 2 O 3 in 10 
 cm. of distilled water in a test tube. Filter if the solution is not clear. 
 Test the action of the nitrate upon litmus paper (?). Conclusions? 
 As 2 O 3 is the anhydride of what acid ? Ascertain the solubility of As 2 O 3 in 
 water by evaporating a portion of the solution to dryness (?). The 
 names of the salts formed by this acid have what ending? Is H 3 AsO 3 
 known in the free state ? What is the best antidote for arsenious oxide ? 
 Write the structural formula for arsenious acid. 
 
 Experiment IV. Salts of Arsenious Acid. Arsenites. 
 
 (a) To a portion of a solution of sodium or potassium arsenite add 
 a little silver nitrate solution (?). Filter. Try the effect of an excess 
 of NH 4 OH upon a portion of the precipitate ( ?). Boil ( ?). What is the 
 effect of HNO 3 upon the precipitate? Equations? 
 
 (b) Add a few cm. 3 of a solution of copper sulphate to a portion of the 
 aqueous solution of the arsenite ( ?). Is the precipitate soluble in NaOH ? 
 Heat the solution (?). Equations? 
 
 Experiment V. Arsenic Acid. Arsenates. 
 
 Note. The student should be on the alert to trace any analogies of 
 crystalline form, composition, solubility, etc., of the phosphates and 
 arsenates. 
 
 (a) Evaporate the solution of arsenic in nitric acid, Exp. I (c\ to dry- 
 ness and redissolve the residue in 20 cm. 3 of hot water. The purpose 
 of this operation is to remove any free HNO 3 . Test the solution with lit- 
 mus paper ( ?). Is the acid reaction due to nitric acid or arsenic acid, 
 H 3 AsO 4 ? Arsenic acid may be prepared by substituting As 2 O 3 for the 
 metallic arsenic and following the foregoing procedure. Equations ? 
 
 (b) Precipitation of silver orthoarsenate. Add a slight excess of Ag- 
 NO 3 solution to 5 cm. 3 of the arsenic acid solution ( ?). Filter. Try the 
 solubility of separate portions of the precipitate in NH 4 OH (?), and in 
 
NITROGEN FAMILY. 221 
 
 HNO 3 (?). What is the effect of boiling the NH 4 OH solution? Equa- 
 tions ? 
 
 (c) Precipitation of magnesium ammonium orthoarsenate. To " mag- 
 nesia mixture" containing an excess of NH 4 OH, add 5 cm. of the solution 
 of H 3 AsO 4 (or any soluble ar senate] ( ?). If the precipitate is not crystal- 
 line dissolve it in the smallest quantity of HC1, and reprecipitate it .with 
 NH 4 OH. Equations? 
 
 (d) Precipitation of ammonium arseno-molybdate. Add an ammoniun 
 molybdate solution to 5 cm. 3 of the H 3 AsO 4 solution (or any soluble arsen- 
 ate) acidified with concentrated HNO 3 ( ?). If precipitation is apparently 
 incomplete, warm the mixture ( ?). Equations ? Name and give the form- 
 ulae of four acids of arsenic. Into what ions does arsenic acid dissociate 
 when placed in water? Is it a "strong" acid? (Refer to tables.) 
 Write the structural formula for the acid. 
 
 Experiment VI. (Optional) Reactions of the Orthoarsenates. 
 
 Exp. VI. "Phosphorus, " may be repeated using any soluble arsenate 
 instead of Na 2 HPO 4 . Equations? 
 
 Experiment VII. Sulphides of Arsenic. Colloidal Solutions. Sulpo-salts. 
 
 (a) Precipitation and properties of arsenic trisulphide. Pass H 2 S into 
 a solution of H 3 AsO 3 (or of an arsenite) (?). What name is applied to 
 this kind of a solution ? 
 
 Pass H 2 S into the solution of As 2 O 3 in HC1 prepared in Exp. Ill (c), 
 (or of an arsenite acidified with HC1) (?). Filter and wash the precipi- 
 tate. Divide the precipitate into two parts. Boil one portion with dilute 
 HC1. Is the precipitate apparently soluble ? Place the other portion of 
 the precipitate in an evaporating dish and add few cm. 3 of warm ammo- 
 nium sulphide, (NH 4 ) 2 S, or warm ammonium polysulphide, (NH 4 ) 2 S X ( ?). 
 Add HC1 to the solution thus formed and note the effect? Equations? 
 
 (b) Precipitation and properties of arsenic pentasulphide. Repeat (a}, 
 substituting H 3 AsO 4 (or an arsenate) for the H 3 AsO 3 . 
 
 Experiment VIII. Tests for Arsenic. 
 
 In addition to Marsh's test, the following tests are used for the detec- 
 tion of arsenic: Reinsch's, Bettendorff's, Gutzeit's, Fleitmann's. Con- 
 sult works on analytical chemistry and make a detailed statement of one 
 of the above methods. Test the method before reporting. The method 
 may be presented in the form of a record of your experiment. 
 
 ANTIMONY, Sb. 
 
 At. Wt. 120. 2 Mol. Wt. (Sb 2 -Sb 4 ). 
 
 The metallic characteristics exhibited by arsenic become more marked 
 with antimony. It is both an acid-forming and base-forming element. 
 
222 EXPERIMENTAL CHEMISTRY. 
 
 It is a weak "metalloid." It gives sets of compounds in which it is triv- 
 alent, and others in which it is quinquivalent. It cannot displace hy- 
 drogen from dilute acids. 
 
 Experiment I. Properties of Antimony. 
 
 Note. The chemical properties of antimony are very similar to those 
 of arsenic. 
 
 (a) Study the physical properties of the elementary substance. Does 
 it possess a metallic appearance ? Is it crystalline ? Heat a small piece 
 of antimony to redness in a hard glass tube closed at one end ( ?). Re- 
 peat the experiment using a tube open at both ends ( ?). 
 
 (b) Heat a small piece of antimony on charcoal in the oxidizing flame. 
 Results ? Will the free element burn in the air ? If so, what is formed ? 
 
 (c) Boil a little powdered antimony in concentrated HNO 3 . Results ? 
 Preserve for use in Exp. IV. 
 
 Hydrogen Derivatives. 
 
 Experiment II. Stibine. Test for Antimony. 
 
 Repeat Marsh's test, using antimony trichloride, SbCl 3 , in HC1 solu- 
 tion in place of AsCl 3 . Distinguish between arsine and stibine. 
 
 Note. The student may be asked to prove experimentally that tartar 
 emetic contains antimony. What is the formula of this salt? 
 
 Sb -f H 3 SbH 3 84,500 cal. 
 Halides. 
 
 Experiment III. Hydrolysis of Antimony Trichloride. Mass Action. 
 
 (a) Recall Exp. II. "Chlorine." Equation? 
 
 (b) Place a small crystal of antimony chloride, SbCl 3 , in a test tube. 
 Add 10 cm. 3 of water and shake (?). Test the liquid \vith litmus paper 
 (?). Prepare a clear solution by adding concentrated HC1, a drop at a 
 time, warming the mixture after each drop. The solution contains what ? 
 To 5 cm. 3 of the solution add a large amount of water. Result ? Explain. 
 What kind of action is this ? Write the equation for the reaction. 
 
 Now add concentrated HC1. Result ? Is the action reversible ? 
 What influences the direction of the reaction ? Write an equation which 
 will represent both actions. Referring to your equation, in which direc- 
 tion does the reaction go with the greatest speed when an excess of acid is 
 used ? When an excess of water is used ? 
 
 (Sb,Cl 3 ) = 91, 400 cal. 
 Oxygen Derivatives. 
 
NITROGEN FAMILY. 223 
 
 Experiment IV. Antimony Trioxide. Antimony Pentoxide. 
 
 Examine the residue which remained in the tube after heating powdered 
 antimony with HNO 3 (Exp. I, (c)). The white insoluble residue is prob- 
 ably a mixture of antimony tri- (Sb 2 O 3 ) and penta- (Sb 2 O 5 ) oxides of anti- 
 mony. These oxides are the anhydrides of what acids? Filter the 
 solution and boil a small portion of the residue with a solution of NaOH 
 until a clear solution is obtained. Products? Does the oxide exhibit 
 acid or basic properties in this reaction ? Equation ? 
 
 Boil another small portion of the residue with strong HC1. Products? 
 Would you infer from this reaction that the oxide of antimony possesses 
 basic properties? Why? Evaporate the clear solution to small bulk. 
 Does it interact with water like SbCl 3 ? Save a portion of the clear solu- 
 tion for Exp. V. 
 
 (Sb 2 , 0,, 3H 2 0) = 167,400 cal. 
 (Sb 2) O 5 , 3H 2 O) = 228,700 cal. 
 
 Experiment V. Sulphides of Antimony. Sulpho-salts. 
 
 Into a solution of SbCl 3 acidified with HC1 pass H 2 S (?). Equation? 
 Filter and wash the precipitate. Divide the precipitate into two parts. 
 To one portion add strong HC1. Warm. Result? Equation? Try 
 the action of warm (NH 4 )S X upon the other portion of the precipitate ( ?). 
 Equation ? To the solution add HC1 ( ?). Is Sb 2 S 3 soluble in (NH 4 ) 2 S X ? 
 What salt is formed? Equations? 
 
 BISMUTH, Bi. 
 
 At. Wt. 208.0 Mol. Wt. (Bi-Bia). 
 
 The metallic character of bismuth considerably exceeds its metalloidal 
 properties. It does not form a hydrogen derivative, and the oxide (Bi 2 O 3 ) 
 which possesses a constitution similar to the acid forming As 2 O 3 exhibits 
 only basic characteristics. Bismuth and its derivatives are usually con- 
 sidered with the metals, but on account of the fact that it forms a number 
 of compounds analogous in composition and properties to the compounds 
 of other members of the group, it is considered in this order. 
 
 Experiment I. Properties of Bismuth. 
 
 (a) Note the physical properties of the elementary substance (?). 
 Is it malleable or brittle. (Test with the pestle.) 
 
 (b) Heat a small piece of bismuth on charcoal in the oxidizing flame. 
 Results ? 
 
 (c) Mix a little oxide or nitrate of bismuth with a small quantity of 
 sodium carbonate, Na 2 CO 3 . Heat the mixture on charcoal in the re- 
 ducing flame. Examine the resulting metallic globule (?). Explain. 
 Equation ? 
 
 (d) Treat a pinch of powdered bismuth with HNO 3 . Result? 
 
224 EXPERIMENTAL CHEMISTRY. 
 
 Products? Concentrate the solution, cool and crystallize. Save for 
 Exp. III. 
 
 Experiment II. Alloys. Wood's Metal. 
 
 Place a small piece of Wood's metal (an alloy) in a small test tube. 
 Support the test tube in a beaker of water; heat the water in the beaker. 
 Determine the temperature at which the alloy melts by taking the tem- 
 perature of the w r ater-bath (?). Name the components of the alloy and 
 give their respective melting points. 
 
 Experiment III. Hydrolysis of Bismuth Nitrate. 
 
 Dissolve the crystals prepared in Exp. I (d), (or crystals of bismuth 
 nitrate, Bi(NO 3 ) 3 ) by heat with the addition of the least possible amount 
 of nitric acid. To a few r cm. 3 of the solution add a large quantity of 
 water. Result? Add a few drops of dilute HNO 3 , just enough to re- 
 move the cloudy appearance (?). Add a large volume of water (?). 
 Explain. Equations ? Is the action reversible ? 
 
 Experiment IV. Preparation of Bismuth Hydroxide. Dehydration. 
 Bismuth Trioxide. 
 
 (a) Try the effect of NH 4 OH upon a clear solution of Bi(NO 3 ) 3 . 
 Result ? What are the products ? Equation ? Filter and wash the 
 precipitate. 
 
 (6) Remove the precipitate from the filter and place it in a porcelain 
 crucible. Ignite. Note the color of the residue of bismuth trioxide 
 when hot; also cold ( ?). Equation ? Define dehydration. 
 
 (Bi 2 ,0 3 ) = ?cal. 
 
 Experiment V. Preparation of Bismuth Trichloride. Hydrolysis. 
 
 Pour a few drops of dilute HC1 upon precipitated Bi(OH) 3 upon a 
 filter (?). Collect the filtrate in a test tube containing 15 to 20 cm. 3 of 
 water. Result? Explain. Equation? The halide of what other element 
 manifests a tendency to undergo hydrolysis ? 
 
 (Bi,Cl 3 ) = 90,600 cal. 
 
 Experiment VI. Bismuth Trisulphide. 
 
 Pass H 2 S through a solution of the chloride or nitrate of bismuth con- 
 taining just enough HC1 to prevent hydrolysis. Result? Products? 
 Equation? Filter and discard the filtrate. Treat the precipitate with 
 warm (NH 4 ) 2 S X . Is the precipitate apparently soluble? Filter and 
 add HC1 to the filtrate. Is the yellowish-white substance free sulphur 
 or bismuth trisulphide, Bi 2 S 3 ? Is Bi 2 S 3 soluble in (NH 4 ) 2 S X ? Equa- 
 tions? Compare results with those in Exp. V. "Antimony" (?). 
 
 The following comparative table will show that the elements of the nitro- 
 
NITROGEN FAMILY. 
 
 22 5 
 
 gen family exhibit the same gradation of properties with increasing atomic 
 weights as was displayed by the elements of the oxygen family: 
 
 Physical Properties Nitrogen 
 
 Atomic weight 14.01 
 
 State or phase Gas 
 
 Color Colorless 
 
 Phosphorus Arsenic Antimony Bismuth 
 
 31.0 75.0 120.2 205.0 
 
 Solid Solid Solid Solid 
 
 (crystalline) (crystalline) 
 
 (crystalline) (crystalline) 
 
 Yellow Steel-gray Silver- Reddish, 
 
 white Metallic. 
 
 Specific gravity 
 Specific heat 
 
 .886 (liq 
 
 uid) 1.82 
 
 5-72 
 
 6.7 
 
 9.8 
 / 
 
 
 
 
 *x 
 
 Melting-point 
 
 2 I 7 
 
 44 
 
 4CO 
 
 42^ 
 
 270 
 
 Boiling-point 
 
 *" O 
 
 -193 
 
 TTT" 
 
 287 
 
 *4-O w 
 800 
 
 t*j 
 
 1,400? 
 
 ^ / w 
 
 i,6oo? 
 
 
 Chemical properties 
 
 
 
 
 
 
 H-derivatives 
 
 NH 3 
 
 PH, 
 
 AsH? 
 
 SbH, 
 
 
 Heat of formation. . 
 
 1 1, 800 
 
 J. -t--l- 3 
 
 cal. 1 1, 600 
 
 1 - OJ -- L 3 ~ "^^j 
 
 cal. 11,700 cal. 84,500 cal 
 
 Chemical properties 
 
 Strong base Weak 
 
 base Neutral 
 
 Neutral 
 
 
 Halides 
 
 NCI, 
 
 PCI, 
 
 AsCl 3 
 
 SbCl 3 
 
 BiCl 3 
 
 Heat of formation . . 
 
 * ^ ^- *3 ~~ v ~o 
 
 38,100 cal. 75,000 cal 
 
 71,400 cal. 
 
 91,400 cal. 
 
 90,600 cal. 
 
 State or phase 
 
 Liquid 
 
 Liquid 
 
 Liquid 
 
 Solid 
 
 Solid 
 
 Stability 
 
 Explosive 
 
 Decomposed 
 
 Decomposed 
 
 Undergoes 
 
 Undergoes 
 
 
 
 by water 
 
 by water 
 
 hydolysis 
 
 hydrolysis. 
 
 O-derivatives 
 
 N 2 3 
 
 P 2 3 
 
 As 2 3 
 
 Sb 2 O 3 
 
 Bi 2 3 
 
 and 
 
 N 2 S 
 
 P 3 0s 
 
 As 2 O s 
 
 Sb 2 O s 
 
 Bi 2 s . 
 
 Chemical properties 
 
 Anhy- 
 
 Anhydrides 
 
 Anhydrides 
 
 Weak anhy- 
 
 Strongly 
 
 
 drides 
 
 
 
 drides, 
 
 basic. 
 
 
 
 
 
 basic. 
 
 
 Acids 
 
 HNO 2 
 
 H 3 PO 3 
 
 H 3 AsO 3 
 
 HSbO 2 
 
 
 IVToin ]]'\r r\ mt ^rrf ioc 
 
 HNO 3 
 
 H 3 P0 4 
 
 H 3 AsO 4 
 
 H 3 SbO 4 
 
 ^ 
 
 PROBLEMS. 
 
 1. If a liter of water absorbs 800 times its volume of ammonia gas, 
 how many grams of ammonia will be absorbed by 500 cm. 3 of water? 
 
 2. How many grams do 100 cm. 3 of concentrated HNO 3 weigh? 
 
 3. How many grams of concentrated HNO 3 can be made from 10 
 grams of sodium nitrate ? 
 
 4. How much phosphorus can be obtained from 100 grams of bone- 
 ash, containing 75 per cent, of calcium phosphate, Ca 3 (PO 4 ) 2 ? From 
 100 grams of pure apatite, Ca 5 F(PO 4 ) 3 ? 
 
 5. Calculate the percentage of arsenic in realgar. In orpiment. 
 
 6. The specific gravity of the vapor of phosphorus relative to hydrogen 
 (H 2 ) is 119.8 at 800. How many atoms in the molecule of the elemen- 
 tary substance ? 
 
 7. Specific gravity of vapor of arsenic at 900 is 296.6; at 1740, 
 154.6. Determine the molecular formula of arsenic at the respective 
 temperatures. 
 
 '5 
 
CHAPTER XX. 
 CARBON FAMILY. 
 
 Carbon C, 12.00 
 
 Silicon Si, 28.4 
 
 (Germanium . . . Ge, 72.5) 
 
 (Tin Sn, 119.0) 
 
 (Lead Pb, 206.9) 
 
 The first two elements of this family are entirely non-metallic, while 
 the others are metals exhibiting properties, however, which reveal 
 resemblances to the non-metals. The group, for advantage of study, 
 may be and is frequently divided into a primary group including carbon 
 and silicon, and a secondary group composed of germanium, tin and lead. 
 All of the elements of the family possess a maximum valence of four. 
 They unite with four atoms of hydrogen or of the halogens. With the 
 exception of silicon, they also form compounds in which they are bivalent. 
 
 PRIMARY GROUP. 
 
 Carbon and silicon resemble one another in their chemical conduct. 
 Some of their physical properties are quite similar. Their derivatives 
 are also very much alike the halogen compounds exhibiting similar 
 properties, the oxides being weak acid anhydrides which are capable 
 of. forming stable salts with bases. The acids of carbon and silicon are 
 unstable, decomposing into water and the acid anhydride. The normal 
 valence of the two elements is four. 
 
 The essential difference in the chemical conduct of carbon and silicon 
 is the ability of atoms of the former to combine with each other and form 
 "chain compounds," i.e., carbon possesses the property of satisfying its 
 own bonds of valence. This property is indicated in the following 
 structural formulae of ethyl alcohol (C 2 H 5 OH), propane (C 3 H 8 ), and 
 butane (C 4 H 10 ): 
 
 H H H H H 
 
 II III 
 
 H C C O H, H C C C H, 
 
 I I I 
 
 H H H H H 
 
 H H H H 
 
 I I I I 
 H C C C C H. 
 
 I I I 
 
 H H H H 
 
 226 
 
CARBON FAMILY. 227 
 
 Silicon does not possess the power of satisfying its own bonds, con- 
 sequently it does not form "chain compounds." 
 
 CARBON, C. 
 
 At. Wt. 12.00 Mol. Wt. (C x ). 
 
 Carbon occurs in the free and almost pure state in nature in several 
 allotropic forms known as diamond, graphite or plumbago and amorphous 
 carbon. It is also found in the combined condition in all living things, 
 in carbon dioxide and carbonates, in turf, peat, lignite or brown coal, 
 bituminous coal and anthracite coal, and in such mineral oils as asphal- 
 tum and petroleum. As the result of the power of carbon to satisfy 
 its own bonds, the number of possible carbon compounds is so large 
 that it appears necessary, for purposes of convenience, to treat them 
 apart in a separate portion or special phase of chemistry. With hydrogen 
 carbon forms practically an unlimited number of compounds into which 
 nearly all other elements, especially oxygen, nitrogen, the halogens and 
 sulphur can enter. The derivatives of carbon have been termed organic 
 compounds because of the idea which prevailed for many years, that 
 these compounds could not be produced without the intervention of 
 life, i.e., that their artificial production was impossible. On account 
 of this original belief the chemistry of the carbon compounds is commonly 
 known as organic chemistry. 
 
 Although it is true tnat a large number of the derivatives of carbon 
 are obtained exclusively from animal and vegetable organisms, yet many 
 of them are prepared artificially from simpler ones or from the elements 
 by simple synthetic methods. The student is also reminded of the fact 
 that hundreds of definite chemical compounds, including drugs and 
 dyes of great value, apparently unknown to either animal or vegetable 
 life, have been prepared artificially. The preparation of many other 
 compounds is delayed solely because of their unusual complexity and 
 instability. 
 
 Experiment I. Properties of Carbon. 
 
 (a) Enumerate the known allotropic modifications of this element. 
 Examine specimens of each. Which represents the purest carbon? 
 Tabulate their principal physical properties ? 
 
 (b) Preparation of amorphous carbon. 
 
 (bi) Place about i gram of cane sugar (C 12 H 22 O n ) in an old test tube 
 and heat until vapors cease to appear. Was water liberated during 
 the process? Pulverize the black residue in a mortar. What is it? 
 Equation ? 
 
 (62) Cover the bottom of a crucible, preferably an iron crucible 
 provided with lid and delivery tube, with sand. Place a number of small 
 pieces of soft pine (largely cellulose, (C 6 H 10 O 5 ) n in the crucible and 
 cover with sand to partially exclude the air. Support the crucible in a 
 
228 EXPERIMENTAL CHEMISTRY. 
 
 pipe-stem or wire triangle and heat for about a half-hour or until the 
 contents of the crucible ceae to smoke. Allow the crucible to cool, 
 then pour the contents out upon an iron plate. Place one or two of the 
 charred pieces of wood in a mortar. It should reduce readily to a fine 
 powder. What is it? Equation? 
 
 (63) Close the holes at the bottom of a lighted Bunsen burner and 
 hold a piece of cold glass tubing in the upper portion of the flame. What 
 is the black deposit ? By what two names is it known ? 
 
 Repeat, using the flame of a candle ( ?). 
 
 (64) (Hood). Apply a flame to a piece of camphor gum (C 10 H 1C O) 
 about the size of a pea. Does it burn with a "sooty" flame? Explain. 
 
 (65) Name the different kinds of charcoal. Give an example of 
 "destructive distillation," and define same. What is boneblack? 
 Coke? 
 
 (c) Porosity of charcoal. Observe that a piece of ordinary charcoal 
 floats upon the water. Place a small piece of it in a test tube half full of 
 water; by means of a long glass rod or a piece of wire force the charcoal 
 down to the bottom of the tube and boil the water for several minutes 
 or until the charcoal shows little or no tendency to float. Explain. 
 
 (d) Decolorizing action of charcoal. To a test tube one-fourth full 
 of powdered animal charcoal add 10 cm. 3 of water which has been 
 colored by the addition of a few drops of an indigo solution or a litmus 
 solution; shake thoroughly for a few minutes, then heat to boiling; filter 
 through a wet filter paper and compare the filtrate with that of the orig- 
 inal solution (?). Explain. 
 
 Note. A solution of brown sugar may be substituted for either of 
 the above solutions. 
 
 (e) Deodorizing action of charcoal. Fill a test tube nearly one-third 
 full of powdered charcoal; add 5 cm. 3 of a weak solution of hydrogen 
 sulphide; cork the tube tightly and shake vigorously; set the tube aside 
 and after fifteen to thirty minutes note the odor of the contents (?). 
 Repeat until your results are definite. Explain. 
 
 (/) Reducing action of charcoal. Mix thoroughly on a piece of paper 
 2 grams of copper oxide, CuO, and .5 grams of powdered charcoal. Place 
 the mixture in a hard glass test tube (ignition tube) ; clamp the tube in a 
 nearly horizontal position and heat strongly. The test tube may be 
 provided with a cork and delivery tube which should dip into a solution 
 of calcium hydroxide, Ca(OH) 2 , in a test tube, or this second test tube 
 containing the Ca(OH) 2 solution may be held vertically under the mouth 
 of the other tube so that the heavy gas which is evolved when the mixture 
 is heated may fall into the tube, then close the tube with the thumb and 
 shake. Explain the milky appearance of the solution. Examine the 
 residue in the ignition tube. This may be done by rubbing it in the 
 mortar and washing away the lighter particles. Results? Explain. 
 Equation ? 
 
 Recall the interaction of As-jOg and charcoal when heated together ( ?). 
 
CARBON FAMILY. 
 
 229 
 
 Equation? What is such a type of reaction called? Is carbon a 
 chemically active substance at ordinary temperatures ? How does the 
 charring of wood preserve it? 
 
 Carbon and Oxygen. 
 
 Experiment II. Preparation and Properties of Carbon Dioxide. 
 
 (a) Recall or repeat Exp. III. "Oxygen" (?). Do carbon and oxygen 
 react at ordinary temperatures? What is formed by their interaction? 
 Equation ? What is the usual test for carbon dioxide ? 
 
 (b) Show that CO 2 k formed when ordinary combustibles, such as 
 wood, paper, illuminating gas, etc., are burned in air. 
 
 FIG. 40. Kipp Generator. 
 
 (c) Hold a piece of charcoal in the Bunsen flame and describe its 
 combustion. Products? Repeat experiment using graphite (pencil 
 lead) (?). Define combustion. 
 
 (d) Pour a few cm. 3 of clear limewater (Ca(OH 2 )) into a test tube 
 and by means of a glass tube blow air from the lungs through the liquid. 
 Result ? Explain. Would you infer that carbon undergoes combustion 
 within the human body? 
 
 (e) Usual laboratory method of preparation (Fig. 40). Put several small 
 pieces of marble (calcium carbonate, CaCO 3 ) in the hydrogen generator 
 and add enough water to seal the bottom of the thistle tube. Connect 
 the generator with a wash bottle half filled with strong sulphuric acid. 
 Pour strong hydrochloric acid into the flask, as needed, through the 
 
230 EXQERIMENTAL CHEMISTRY. 
 
 thistle tube. Collect five or six bottles of the gas by displacement of 
 air, placing the bottle with the mouth upward. Prove that the evolved 
 gas is carbon dioxide. Note its color, odor and taste. Is the gas com- 
 bustible? Does it support combustion? Pour one bottleful of the gas 
 very slowly into another bottle of about the same size containing air. 
 Lower a lighted splinter into the second bottle. Result ? Conclusion ? 
 Is the gas heavier or lighter than air ? 
 
 Counterpoise a beaker on the balance, then pour carbon dioxide 
 into it (?). 
 
 Filter the contents of the generating flask; evaporate (hood) a portion 
 of the filtrate to dryness; heat the residue until fumes of hydrochloric 
 acid cease to be evolved. Dissolve a portion of the residue in distilled 
 water and test for a chloride. Test for calcium by dipping a clean 
 moist platinum wire into the solid residue and holding it in the Bunsen 
 flame. A yellowish-red flame indicates the presence of calcium. What 
 is the name of the salt (residue) in the evaporating dish? Indicate by 
 equations the reactions involved in the preparation of carbon dioxide. 
 
 (/) Try the action of hydrochloric acid upon small portions of each 
 of the following salts: sodium carbonate (Na^COJ, sodium hydrogen 
 carbonate (NaHCO 3 ), potassium carbonate (K 2 CO 3 ), barium carbonate 
 (BaCO 3 ). Results? Is CO 2 evolved in each case? (Use the "loop 
 tube " to carry a film of lime-water to the mouth of the test tube. This 
 renders the testing for carbon dioxide a very simple process.) 
 
 (g) Place about a gram of sodium carbonate or powdered magnesite 
 (MgCO 3 ) in a test tube, and heat strongly. Is carbon dioxide evolved? 
 
 Interpret the following equations: 
 
 C (charcoal) + O 2 > CO 2 + 96,980 cal. 
 
 C (retort carbon) + O 2 > CO 2 + 96,530 cal. 
 
 C (graphite) + O 2 CO 2 + 93,360 cal. 
 
 C (diamond) + O 2 > CO 2 + 93,240 cal. 
 
 Experiment III. Optional. Synthesis of Carbon Dioxide. 
 
 This operation may be carried out by drawing simultaneously purified 
 air through a tube over hot charcoal, and the product into a solution of 
 calcium or barium hydroxide.. (Instructions.) 
 
 Experiment IV. Carbonic Acid. Carbonates. 
 
 (a) Half fill a clean bottle with distilled water; test the water with 
 litmus paper or a few drops of a solution of phenolphthalein ( ?). Allow 
 a stream of carbon dioxide to bubble through the water until the latter 
 is saturated. Test the water with litmus paper. Results? Con- 
 clusions? Ascertain the relative conductivity of the solution (?). CO 2 
 is the anhydride of what acid? (This acid, carbonic (H 2 CO 3 ), has 
 never been prepared in the free state.) Is the acid stable? If not, 
 into what substances does it decompose ? What ions does the acid yield ? 
 Is it a " strong" acid ? State the reasons for your answer. 
 
CARBON FAMILY. 231 
 
 Can you explain why carbonic acid is so readily liberated from car- 
 bonates? State Berthollet's law. Henry's law. 
 
 (CO 2l Aq.) = 5800 cal. 
 
 (b) Conduct carbon dioxide into lime-water, or baryta water (Ba(OH) 2 ) 
 until the gas ceases to be absorbed. Results? Filter and test the 
 precipitate for carbon dioxide (?). What salt is formed by the action 
 of CO 2 on lime-water? Is it soluble in water? In acids? Equations? 
 
 Saturate a solution of sodium hydroxide with carbon dioxide. Set 
 the solution aside and allow the liquid to evaporate spontaneously. 
 Prove that the residue is a carbonate. Equations ? 
 
 What class of salts are yielded by carbonic acid (H 2 O + CO 2 ) ? What 
 is the basicity of carbonic acid? How would you determine whether a 
 rock specimen contained a carbonate or not ? 
 
 Experiment V. Formation of Calcium Acid Carbonate. 
 
 Conduct carbon dioxide into lime-water until the precipitate which 
 forms at first, disappears. Filter. Divide the filtrate into two portions. 
 To one portion add clear lime-water (?). Heat the second portion in 
 a test tube ( ?). Why did the precipitate disappear in the first reaction ? 
 Why does it appear after heating ? Equations ? Explain the formation 
 of the incrustations on the inside of tea-kettles and steam boilers. 
 
 All normal carbonates except those of the alkalies are insoluble in 
 water. They are decomposed by hydrochloric acid with an evolution 
 of carbon dioxide. 
 
 Experiment VI. Optional. (Quant.) Estimation of Carbon Dioxide in 
 a Carbonate. 
 
 (a) Report a method for the determination of carbon dioxide in a 
 carbonate in which the evolved gas is absorbed in a weighed apparatus. 
 
 (b) Estimation by difference. The epitome of the process is as 
 follows: A known weight of the carbonate is decomposed by dilute 
 acid in a weighed apparatus and the carbon dioxide is dried as it escapes 
 or is expelled through a bulb containing strong sulphuric acid. The 
 loss in weight which results represents the carbon dioxide in the sub- 
 stance being examined. 
 
 Schrotter's apparatus is a convenient form of the apparatus required 
 for the performance of this experiment. Draw a diagram of the appa- 
 ratus and explain its manipulation. (See model in laboratory or text 
 books on analytical chemistry. Figures of various forms of the apparatus 
 usually appear in catalogues of laboratory supplies.) 
 
 Experiment VII. (Quant.) Determination of the Density of Carbon 
 Dioxide. Molecular Weight. 
 
 Provide a 300 cm. 3 flask with a tightly fitting rubber cork. See that 
 the apparatus is perfectly clean and dry. Weigh the flask and cork. 
 
232 EXPERIMENTAL CHEMISTRY. 
 
 Record this weight as the weight of flask, cork, and air which filled flask. 
 Fill the flask with carbon dioxide by displacement of air; stopper the 
 flask and weigh. Record weight. Repeat the operation of filling the 
 flask with carbon dioxide and weighing, until the weight becomes practi- 
 cally constant. Calibrate the flask as in previous experiments. The 
 weight of the empty flask is found by subtracting from the weight of the 
 cork and flask filled with air the weight of a volume of air equal to the 
 capacity of the flask. (Air under normal conditions, i cm. 3 = .00129 
 gram.) What is the weight of the carbon dioxide in the flask? Of i 
 litre of the gas? 
 
 Calculate the molecular weight of carbon dioxide from the experi- 
 mental data (?). What is your percentage of error? 
 
 Mention sources of error in determining the density of the gas by above 
 method. 
 
 Experiment VIII. Preparation and Properties of Carbon Monoxide. 
 
 (a) Assemble apparatus similar to that which you used for the pre- 
 paration of chlorine. Put 10 grams of oxalic acid (C 2 H 2 O 4 ) in the gener- 
 ating flask. Half fill the wash bottle with lime-water or a caustic-soda 
 solution. Pour 35 cm. 3 of concentrated sulphuric acid through the 
 thistle tube into the flask; heat the latter gently. Collect three or four 
 bottles of the evolved gas over water. Add a few drops of lime-water 
 to a bottle of the gas; place your hand over the mouth of the bottle and 
 shake vigorously. Is there any evidence that carbon dioxide is present ? 
 Set fire to a bottle of the gas and notice the characteristic blue flame. 
 After the gas has burned, pour a few cm. 3 of lime-water into the bottle; 
 shake. Results ? Describe any visible changes that may have occurred 
 in the wash-bottle. Add a few drops of acid to the solution and test 
 for carbon dioxide (?). What gas besides carbon monoxide was evi- 
 dently evolved by the action of H 2 SO 4 on C 2 H 2 O 4 ? This chemical 
 action is due largely to what particular property of sulphuric acid? 
 Equations? May CO act as a reducing agent? Explain. Interpret the 
 following equations (?). 
 
 (b) Optional. Intimately mix 2 grams of zinc oxide (ZnO) and i 
 gram of powdered charcoal. Introduce the mixture into a hard glass 
 test tube and heat strongly. Apply a flame to the mouth of the tube. 
 Result ? Equations ? 
 
 (C, O) = 28,500 cal. (CO,O) = 68,400 cal. 
 
 Cu,O (37,100 cal) + C,O (28,500 cal.) > Cu + C,O 2 (96,900 cal.) 
 + 31,300 cal. 
 
 " If an element combines with another according to multiple pro- 
 portions, there usually occurs, in the union of the first atom, a greater 
 disengagement of heat than with the following atom (compare nitrogen 
 oxides). The numbers above, on the contrary, show that the union 
 of the second atom of oxygen wuth carbon (CO,O) sets free 68,400 
 calories; that of the first atom (C,O), however, only 28,500 calories. 
 
CARBON FAMILY. 233 
 
 This can only be explained by the fact that, for the vaporization and 
 disaggregation of the solid carbon molecules, heat is necessary. If we 
 assume that the direct union of the first atom also disengaged 68,400 
 calories, it would follow from this that, in the dislocation of twelve 
 parts carbon (amorphous) by weight into gaseous free atoms, 39,900 
 (68,400-28,500) calories were absorbed." Richter. 
 
 The above equations, indicating the heat disengagement for each 
 stage in the oxidation of carbon, furnish an excellent illustration of the 
 Law of Hess. (State the law.) 
 
 Experiment IX. (Quant.) Determination of the Density of Carbon 
 Monoxide. Molecular Weight. 
 
 The density of carbon monoxide may be determined by a method 
 identical with that recorded in Exp. III. " Nitrogen and the Atmosphere.'' 
 
 Carbon and Hydrogen (Hydrocarbons). 
 
 Experiment X. Preparation and Properties of Methane (Marsh Gas). 
 
 Heat a few grams of sodium acetate in a porcelain dish until the water 
 of crystallization has been expelled. Heat a mixture consisting of 
 equal parts of dehydrated sodium acetate and soda lime in a hard glass 
 test tube clamped in a nearly horizontal position. The tube should be 
 fitted with a cork and delivery tube. Collect the gas over water. Note its 
 odor, color and taste. Does it burn ? Does the gas explode when mixed 
 with air ? Attach a jet to the delivery tube and burn the gas when all the 
 air has been displaced from the apparatus. Note the degree of luminosity 
 of the flame. Is water formed during the burning of the gas? Test. 
 Carbon dioxide ? Test. Equations ? What is the name and molecular 
 formula of the gas? 
 
 (C,* H 4 ) = 21,700 cal. (C 2 , H 6 ) = 28,600 cal. 
 
 Experiment XI. (Optional.) Preparation of Ethylene by Elimination 
 of Water from Alcohol. Ethylene Bromide. 
 
 The following method is suggested by Gattermann: A mixture of 
 25 grams of alcohol and 150 grams of concentrated H 2 SO 4 is heated, 
 not too strongly, in a liter round flask on a wire gauze covered with thin 
 asbestos, or sand bath. As soon as an active evolution of ethylene 
 takes place, add, through a dropping funnel, a mixture of one part 
 alcohol and two parts concentrated H 2 SO 4 (made by pouring 300 grams 
 of alcohol into 600 grams of sulphuric acid, with constant stirring), 
 slowly, so that a regular stream of gas is evolved. If the mixture in the 
 flask forms badly with a separation of carbon, it has been too strongly 
 heated, and it is advisable to empty the flask and begin the operation 
 anew. In order to free ethylene from alcohol, ether and sulphur dioxide, 
 it is passed through a wash bottle containing sulphuric acid, and a second 
 
 # Amorphous carbon. 
 
234 
 
 EXPERIMENTAL CHEMISTRY. 
 
 one, provided with three tubulures, the central one supplied with a safety- 
 tube, containing a dilute solution of caustic soda. It is well to use a 
 third wash bottle containing sulphuric acid. When the air in the 
 apparatus has been displaced, fill a small narrow-mouthed bottle provided 
 with a greased glass stopper, with the gas by downward displacement 
 of air. Allow a drop of bromine to fall into the bottle, then replace the 
 stopper quickly. Observe any changes taking place within the bottle 
 After a few minutes remove the stopper under water. Results ? Explain. 
 Attach a jet to the delivery tube and burn the gas. Compare the 
 degree of luminosity of the flame with that of methane (?). Equations? 
 
 (C 2 , H 4 ) = - 2700 cal. 
 Experiment XII. Preparation and Properties of Acetylene. 
 
 Fit a test tube with a single-hole stopper carrying a piece of straight 
 glass tubing about 5 cm. long and ending flush with the lower side of the 
 stopper. Fill the test tube nearly full of water then drop into it a small 
 piece of calcium carbide, CaC 2 ; insert the stopper. Note the odor of 
 
 FIG. 41. The Manufacture of Coal Gas. 
 
 the evolved gas. When the air in the tube has been displaced, light 
 the gas which is issuing from the jet. Describe the character of the 
 flame. Hold a glass plate just above the flame (?). Would you infer 
 that there is perfect combustion of carbon ? Why ? What is the name 
 and molecular formula of the gas? Equations indicating the action of 
 the carbide on water and the combustion of the gas? Will air burn in 
 an atmosphere of the gas?' Can you offer a possible explanation as to 
 
CARBON FAMILY. 235 
 
 why C 2 H 2 and C 2 H 4 are endothermic compounds, and CH 4 and C 2 H 6 
 
 are exothermic? (See carbon monoxide.) (C 2 H 2 ) = 47,800 cal. 
 
 Coal Gas is formed by the destructive distillation (Fig. 41) of bituminous 
 coal. The composition of the products obtained depend largely upon the 
 nature of the coal and the process of distillation. These products are: (i) 
 coke; (2) coal tar; (3) gas liquor containing ammonia and other products, 
 and known as ammoniacal liquor; and (4) coal gas. The coal gas is 
 a mixture of various gases. The components of the mixture may be 
 divided into three classes as follows: 
 
 Illuminants. Diluents. Impurities. 
 
 (About 6%) (About 90%) (About 3%) 
 
 Ethylene, C 2 H 4 Hydrogen, H 2 Nitrogen, N 2 
 
 Propylene, C 3 H e Methane, CH 4 Carbon dioxide CO 2 
 
 Butylene C 4 H 8 Carbon monoxide CO Hydrogen sulphide, H 2 S 
 Acetylene C 2 H 2 Ammonia, NH 3 
 
 Allylene, C 3 H 4 
 Benzene, C 6 H 
 Heavy hydrocarbons. 
 
 Some of the impurities are removed entirely from the purified gas; 
 i.e., before the gas is delivered to the gas mains. The gas is often called 
 "illuminating gas." The student is referred to Sadtler's Hand-book 
 of Industrial Organic Chemistry and Thorpe's Outlines of Industrial 
 Chemistry. 
 
 Experiment XIII. Preparation and Properties of Illuminating Gas. 
 
 Fill a hard glass test tube one-half full of coarsely powdered bituminous 
 or cannel coal; place a plug of glass wool or shredded asbestos above 
 the coal to hold it in place; clamp the tube in a horizontal position and 
 connect it with an empty wash bottle which communicates with a U-tube 
 containing strips of litmus paper in one limb and filter paper moist- 
 ened with lead acetate in the other limb; a third wash bottle is one-third 
 full of lime-water. Heat the ignition tube gently with the Bunsen flame. 
 As soon as all of the air has' been driven out of the apparatus collect two 
 bottles of the gas by water displacement. Disconnect the apparatus. 
 Examine the gas in the bottles, noting its color, odor, inflammability 
 and color of flame. Of what does the gas consist ? Does the residue 
 in the ignition tube suggest some form of carbon ? What name is applied 
 to it ? Describe in full and explain any changes which may have occurred 
 within the wash-bottles. What are the products of the combustion of 
 the chief components of illuminating gas? Devise an experiment to 
 prove your conclusions. (Hint. Use gas from jet.) 
 
 Make a brief statement of the composition and manufacture of 
 "producer gas." Of "water gas." 
 
 Carbon and Sulphur. 
 
236 EXPERIMENTAL CHEMISTRY. 
 
 Experiment XIV. Properties of Carbon Bisulphide. 
 
 Procure a few cm. 3 of carbon disulphide in a test tube from the side- 
 shelf reagent. Note its odor, color and high refractive index. What 
 is its specific gravity ? Place a drop of the liquid on the hand and force 
 air over it ( ?). Place a very small quantity of the substance in an evapo- 
 rating dish, then bring the heated end of a glass rod near the surface of 
 the liquid (?). Would you infer that carbon disulphide is very inflam- 
 mable ? Is it miscible with water ? With alcohol ? It has been used in 
 previous experiments as a solvent for what substances? Is rubber 
 soluble in it ? If so, allow the liquid to evaporate spontaneously. Result ? 
 Enumerate other uses of carbon disulphide. What is the molecular 
 formula of carbon disulphide ? Is it an endothermic compound ? (See 
 following equation.) What would you infer as to its stability? 
 
 (C, S 2 liq.) = 19,600 cal. 
 
 Carbon and the Halogens. Although carbon does not combine di- 
 rectly with the members of the halogen family, halides of carbon are 
 formed by the action of the halogens on the hydrocarbons. (Recall Exp. 
 XL) In our previous work we have observed that chlorine is capable 
 of withdrawing or substituting itself for the hydrogen of water, hydrogen 
 sulphide, ammonia, etc. The halogens react very similarly with the 
 hydrocarbons forming, by a process termed substitution, a series of 
 carbon compounds known as substitution products. By such reactions 
 the following familiar substances are obtained: Methyl-chloride (CH 3 C1), 
 chloroform (CHC1 3 ), iodoform (CHI 3 ), carbon tetrachloride (CC1 4 ), 
 etc. The latter (CC1 4 ) is a colorless liquid having a peculiar odor. 
 It has a sp. gr. of 1.6 and boils at 77. It is the final product of the 
 action of chlorine upon methane (CH 4 ) or upon chloroform. 
 
 CH 4 + C1 2 CH 3 C1 + HC1 
 CH 4 + 2 C1 2 CH 2 C1 2 + 2HC1 
 
 CH 4 + 4 C1 2 CC1 4 + 4HC1 
 
 (C, C1 4 gas) = 21,000 cal. (C, C1 4 liq.) = 28,300 cal. 
 (C 2 , C1 4 gas) = 1 100 cal. 
 
 The heats of formation of the chlorides of carbon approximate those 
 of the hydrogen derivatives, indicating that the affinity of the two ele- 
 ments for carbon is practically the same. 
 
 Carbon and Nitrogen. These two elementary substances apparently 
 do not possess any direct affinity for one another, therefore the union 
 of the two is usually effected by indirect processes. However, small 
 quantities of cyanogen (CN or Cy) are formed when electrical dis- 
 charges take place between carbon poles in an atmosphere of nitrogen. 
 The production is facilitated if one, at least, of the elements is in the 
 form of a compound. Cyanogen is an endothermic compound a 
 colorless gas with an odor resembling that of peach blossoms, or better, 
 
CARBON FAMILY. 237 
 
 almonds. It burns with a characteristic violet-blue flame, or more 
 accurately, with a "pink flame edged with green." Its sp. gr. is 1.805 
 (Air = i). Nascent cyanogen is recognized as a radical or group with 
 the formula, CN. It is seen at once from its structure ( C EEN) that 
 it is an unsaturated compound a negative univalent radical. Like 
 other univalent groups (OH, NH 2 , CH 3 ), it cannot exist in the free state, 
 but combines with itself to form a double molecule of free cyanogen (C 2 N 2 ), 
 known as dicyanogen. The structural formula (N == C C EE N) 
 shows carbon satisfying its own bonds. 
 
 In its chemical relations carbon exhibits the mutability peculiar to 
 organic groups. It combines with hydrogen under the,, influence of a 
 quiet electric discharge to form the very "weak" but extremely poisonous 
 hydrocyanic or Prussic acid (HCN). This acid is a colorless mobile 
 liquid which boils at about 27. 
 
 An aqueous solution possesses small electrical conductivity. It forms 
 cyanides by exchanging its hydrogen for metals. These salts, with the 
 exception of potassium cyanide which undergoes hydrolysis to a marked 
 degree, are relatively stable and uniformly poisonous. The extreme 
 feebleness of the acid is shown by the fact that the carbon dioxide of the 
 air liberates it. For this reason the salt usually possesses the odor of 
 hydrocyanic acid. 
 
 The tendency of the salts, like potassium cyanide (KCN), in the 
 presence of easily reducible and highly heated oxides to form cyanates 
 (KCNO) indicates that hydrocyanic acid is an unsaturated compound, 
 and possesses a structure similar to that indicated by the following 
 formula, H N = C = . This behavior of potassium cyanide accounts 
 for its frequent use as a suitable flux for securing reductions. It mani- 
 fests a similar tendency to combine with sulphur, forming a salt belonging 
 to a class known as the thiocyanates, e.g., KCNS. 
 
 Cyanogen also unites with iron to form a compound or rather a group 
 which combines. with various other metals to produce two series of salts, 
 namely, the jerrocyanides and the jerricyanides. The potassium salt 
 of the former, potassium ferrocyanide (K 4 Fe(CN) 6 .3Aq.), has been 
 given the name of yellow prussiate oj potash; the corresponding salt of 
 the latter is known as potassium ferricyanide, K 3 Fe (CN) 6 or K 6 FeCy 6 . 
 
 (C, N, H gas) = -27,500 cal. 
 Experiment XV. Preparation and Properties of Cyanogen. 
 
 Note. This gas is intensely poisonous and should not be prepared 
 by the student. 
 
 Cyanogen may be prepared by (i) heating the cyanides of mercury 
 or silver; (2) the dry distillation of ammonium oxalate, and (3) the fusion 
 of potassium cyanide and mercuric chloride. 
 
 Hg(CN) 2 Hg + "QN L 
 (NH 4 ) 2 C 2 4 -> 4 H 2 + C 2 N 2 
 2 KCN + HgCl 2 Hg + 2 KC1 + C 2 N 2 
 
238 EXPERIMENTAL CHEMISTRY. 
 
 Experiment XVI. Preparation and Properties of Potassium Cyanide. 
 Potassium Thiocyanate. 
 
 (a) Potassium cyanide (KCN) may be prepared by heating a nitrogen- 
 ous carbon compound like gelatin (C^H^N^O^, approx.), with a 
 small piece of potassium. The KCN which forms is extracted with 
 water. 
 
 To a few cm. 3 of a silver nitrate solution add a few drops of a solution 
 of KCN. Results ? Add an excess of KCN. Result ? Equations ? 
 
 (6) A test for cyanogen compounds. To a dilute aqueous solution of 
 KCN add equal quantities of a ferrous and a ferric salt. Result ? Add 
 an excess of a potassium hydroxide solution, then acidify with HC1. 
 Results? Before attempting to state the composition of the final pro- 
 duct, add a few drops of potassium ferrocyanide (K 4 FeC 6 N 6 ) to a solution 
 of a ferric salt. Result? What are your conclusions now as to the 
 composition of the product referred to above ? 
 
 (c) Potassium thiocyanate. Place two or three cm. 3 of a KCN solu- 
 tion in an evaporating dish and add ammonium polysulphide (NH 4 ) 2 S X 
 until the color of the latter persists. Evaporate to dryness. Dissolve 
 a portion of the residue in water; add a few drops of a ferric chloride 
 (FeCl 3 ) solution (?). Test a solution of potassium thiocyanate (KCNS), 
 shelf reagent, with a few drops of FeCl 3 ( ?). If the solution you prepared 
 gave a black precipitate when FeCl 3 was added, heat the residue again 
 in the evaporating dish and repeat test with FeCl 3 . What is the name 
 of the prepared substance ? Its formula ? Equations ? What is the 
 action of a solution of potassium thiocyanate upon solutions of ferrous 
 salts ? 
 
 Test the tap water for iron by evaporating 100 cm. 3 to small bulk; 
 add a few drops of HNO 3 (?) and heat to boiling. Cool the solution, 
 and add a few drops of a KCNS solution (?). 
 
 All the single cyanides, except those of the alkalies, alkaline earths 
 and mercury (ic) are insoluble in water. Barium cyanide is diffi- 
 cultly soluble. 
 
 Experiment XVII. A Study of Flames. 
 
 (a) Recall or repeat the experiment with the Bunsen burner. 
 Preliminary Exercises. 
 
 (b) Light a Bunsen burner. Bring a cold piece of brass or iron wire 
 gauze down upon the flame. Observe that the flame does not pass 
 through the gauze (Fig. 42). Is this due. to the absence of combustible 
 gases? 'Apply a lighted match above the gauze (?). Why did the flame 
 not pass through the gauze ? Turn off the gas; hold a piece of gauze 3 or 
 4 cm. above the top of the burner; turn on the gas and hold a lighted 
 match above the gauze (?). Explain results. Define "kindling tem- 
 perature." What is the principle upon which the miner's safety lamp 
 (Fig 43) is constructed? 
 
CARBON FAMILY. 239 
 
 (c) Open and close the holes of the lighted burner several times. 
 Results? Hold a piece of glass tubing in the "luminous flame" of the 
 burner (?). Open the holes at the bottom of the burner and hold the 
 blackened rod in the non-luminous flame ( ?). Account for the deposition 
 of carbon and its disappearance. What is the probable cause of the 
 luminosity of the flame? Test your conclusion as follows: Place a 
 small quantity of powdered charcoal in a piece of glass tubing and blow 
 the particles of carbon into one of the holes at the base of the burner. 
 What is the effect on the luminosity of the flame ? Explain. Why do a 
 number of the "burning oils" burn with a "sooty" flame? Why is the 
 hydrogen flame non-luminous? What is the source of the powerful 
 light emitted by a Welsbach burner ? What is the object of the holes at 
 the base of the burner ? Why is the Bunsen flame non-luminous? Why 
 
 FIG. 42. FIG. 43. Davy's Safety Lamp. 
 
 does the gas ordinarily burn at the top of the tube and not on the inside ? 
 What causes the flame to "strike back" occasionally and burn inside 
 the tube? 
 
 (d) Slightly bend a glass tube about 15 cm. long; introduce the shorter 
 arm in the flame about 2 cm. above the top of the burner. Light the gas 
 issuing from the tube. What are your conclusions concerning the con- 
 ditions existing in this portion of the flame ? Explain. 
 
 (e) Oxidizing and reducing flames. Examine and sketch the parts 
 of a very small luminous flame. 
 
 Hold a piece of bright copper wire horizontally across a Bunsen flame 
 so that the wire cuts the inner cone. Observe that the portion of the wire 
 in the inner cone remains bright, while those portions in contact with the 
 edges of the flame become coated with a dark substance (copper oxide, 
 CuO). The copper has undergone what kind of a chemical change? 
 Move the wire so as to bring the oxidized portion into the inner cone, 
 the metal becomes bright owing to a reduction of the oxide. The outer 
 
240 
 
 EXPERIMENTAL CHEMISTRY. 
 
 cone of the flame where oxygen is in excess is called the oxidizing flame ; 
 the inner cone, in which heated and unburnt combustible gases exist 
 (hydrogen and hydrocarbons), is referred to as the reducing 
 flame. Both cones exist in all ordinary flames. 
 
 (/) Use of the blow-pipe. Ask the instructor how to 
 produce the oxidizing and reducing flames by means of a 
 blow-pipe. Ascertain the effect of each flame by heating an 
 intimate mixture of lead oxide (PbO) and sodium carbonate 
 on charcoal in the reducing cone, and a small piece of 
 metallic lead on charcoal in the oxidizing cone. Results ? 
 
 (g) Examine and sketch a candle flame (Fig. 44). Is there 
 any essential difference between a gas flame and a candle 
 flame or a lamp flame ? All flames are the result of the 
 interaction of what state of matter? Define "a flame." 
 
 (h) Examine a lamp burner. Does its structure embody 
 principles analogous to those of the Bunsen burner? Ex- 
 plain. Why are some lamps provided with a "central 
 draft?" Can you explain the cone-like shape of the flame 
 FIG. 44. o f t^ Bunsen burner ? 
 
 The following table gives the results of a series of accurate experiments 
 conducted for the purpose of determining the actual temperatures of 
 various regions of the Bunsen flame. The cooling effect of the diluent, 
 air, is observed in portions of the non-luminous flame. 
 
 TEMPERATURE OF FLAME OF BUNSEN BURNER. 
 (Data for a burner burning six cubic feet of coal gas per hour.) 
 
 (Lewes-Newth.) 
 
 Region in Flame. Luminous. 
 
 One-half inch above burner I 35- 
 
 One and one-half inch above burner . . 421. ... 
 
 Tip of inner cone 913 . . . . 
 
 Center of outer cone 1328. ... 
 
 Tip of outer cone 728. ... 
 
 Side of outer cone, level with tip of in- 
 ner cone 1236. . . . 
 
 Non-luminous. 
 54 
 
 1090' 
 
 1333' 
 
 SILICON, SI. 
 
 At. Wt. 28.4 Mol. Wt. ( ? ). 
 
 It has been estimated that 27.3 per cent, of the earth consists of silicon. 
 However, owing to its great affinity for oxygen, it does not occur in 
 nature as an uncombined element. In the combined condition it most 
 frequently occurs as silicon dioxide (silica, SiO 2 ) and in the form of salts 
 of silicic acid silicates. Silicon, like carbon, occurs in several forms, 
 
CARBON FAMILY. 241 
 
 of which amorphous and crystalline forms are accurately known. 
 Amorphous silicon is a non-lustrous greenish-brown powder which when 
 set on fire in the air undergoes incomplete combustion owing to the 
 formation of an incombustible layer of silicon dioxide. 
 
 Crystalline silicon (sp. gr., 2.49) is a grayish-black octohedral crystal 
 possessing a metallic luster. It does not change appreciably in the air 
 even at red heat. It is soluble in hot caustic soda. (Molten iron or 
 zinc dissolves silicon; when they cool, the silicon separates in crystals.) 
 
 The most important compound of silicon is silicon dioxide, the anhy- 
 dride of silicic acid. It occurs in several v arieties two crystalline and 
 one amorphous. It is most widely distributed in the crystalline form as 
 quartz or rock crystal. There are several colored varieties of quartz, 
 the color being due to the presence of impurities. Quartz has a density 
 of 2.66 and stands seventh in the scale of hardness. (See Appendix.) 
 The other crystalline form of silicon dioxide is known as tridymite. It 
 occurs almost solely in the form of microscopic crystals as a constituent 
 of various rocks. The sp. gr. of tridymite is 2.3. Amorphous silicon 
 dioxide (sp. gr., 2.2) occurs most abundantly in the form commonly 
 known as flint. The opal represents another impure form. Jasper, 
 chalcedony, etc., are "cryp to-crystalline," i.e., are composed of small 
 crystals. Silica dissolves in the fused hydroxides and carbonates of the 
 alkalies, but is insoluble in water and all acids except hydrofluoric acid. 
 At the high temperature of the oxyhydrogen flame silicon dioxide melts 
 forming a viscous glass-like fluid which can be blown into any desired 
 shape. On account of its high melting-point and resistance to chemical 
 action, this amorphous "quartz-glass" has been used extensively in 
 recent years for making vessels for use in chemical laboratories. It 
 has a very small coefficient of expansion with heat and can, therefore, 
 be subjected to sudden changes of temperature without cracking. 
 
 In its chemical relations silicon bears a close resemblance to carbon. 
 It is essentially non-metallic in these relations. The natural silicates 
 form a numerous and complex class of minerals which may be regarded 
 as compounds of metallic oxides with silicon dioxide; as feldspar, A1 2 O 3 ,- 
 K 2 O, 6SiO 2 . 
 
 Experiment I. Preparation and Properties of Silicon. 
 
 (a) Procure a piece of crystalline silicon from the instructor. Ex- 
 amine it carefully. Tabulate its observed properties. What is its sp. gr. ? 
 Enumerate the several forms in which silicon occurs. 
 
 (6) Thoroughly mix in a mortar i gram of magnesium powder and 2 or 
 3 grams of powdered quartz sand. Place the mixture in a wide hard glass 
 tube; heat the tube to bright redness in the flame of a blast-lamp, slowly 
 rotating the tube. After a few minutes' heating, dissolve out the soluble 
 portions of the fused mass with water and HC1. The residue consists 
 of amorphous silicon and quartz sand. 
 
 (c) Ascertain the action of the following reagents upon silicon and 
 tabulate results: HC1, H 2 SO 4 , HNO 3 , and a KOH solution. 
 16 
 
242 EXPERIMENTAL CHEMISTRY. 
 
 Experiment II. Silicon Dioxide. Silicates. Silicic Acid. Hydrogele. 
 
 (a) Examine specimens of quartz, sand, agate, opal, amethyst, onyx, 
 flint, infusorial or diatomaceous earth, etc. Make a list of those 
 examined. 
 
 (b) Test the solubility of silicon dioxide in the various acids and 
 alkalies. Tabulate results. Recall the action of H 2 F 2 on silica (SiO 2 ). 
 See Exp. Ill, "Fluorine." Equation ? 
 
 (c) To 2 cm. 3 of a solution of sodium silicate (water glass, Na 2 SiO 3 ) 
 in an evaporating dish or casserole, add 1 5 or 20 cm. 3 of water, then add con- 
 centrated HC1 until the solution is strongly acid. Note the formation of a 
 white gelatinous precipitate. What is it ? Formula ? Equation ? Note. 
 These jelly-like precipitates are sometimes termed " hydro geles." 
 
 Evaporate the solution to dryness on a sand or steam bath. Now 
 heat the dish with the flame. When the dish has cooled, treat the residue 
 with water; filter. Remove the residue from the filter paper and heat it 
 in a covered crucible supported by a pipe-stem triangle. Results ? What 
 is the substance in the crucible ? Its formula ? Rub some of it between 
 the fingers. Do you detect any grit ? Is it crystalline or amorphous ? 
 
 Sil-ic acid (probably H 4 SiO 4 ) shows a tendency to form "colloidal 
 solutions" or "pseudo-solutions." When hydrochloric acid is added 
 to a concentrated solution of a soluble silicate, the silicic acid separates 
 out as a gelatinous mass, but if a dilute solution of the silicate is sub- 
 stituted and an excess of the acid used, precipitation does not take place 
 and the solution remains clear and apparently unchanged. It appears 
 as if the silicic acid were difficultly soluble in water, but dissolved when 
 much water was used. It has been found, however, that this is not the 
 case. The silicic acid is present (suspended) in the water in the colloidal 
 state. The apparent solution of silicic acid is not a true solution, but a 
 "pseudo-solution." (A solution of this kind is sometimes termed a 
 hydrosole.) This may be proven by subjecting the pseudo-solution to 
 dialysis, when the salt (crystalloid) formed by the interaction of the 
 metal and acid will diffuse freely through the membrane of a dialyser, 
 while the silicic acid will be retained, as is in accord with the conduct 
 of colloidal substances. Again, it is possible to partially precipitate the 
 silicic acid from these "pseudo-solutions" by the addition of various 
 substances, such as salts. Silic acid occurs very frequently in this form 
 in nature as the result of the decomposition of certain silicates by carbonic 
 acid. The various colored varieties of quartz were probably formed 
 from silicic acid in this condition. This view is supported by the fact 
 that it is now known that silicic acid, under suitable conditions, will 
 crystallize out from such solutions. 
 
 (b) Fuse a mixture of i grm. of finely powdered feldspar with 4 or 5 
 grams of sodium carbonate* in a covered porcelain or platinum crucible. 
 
 * A mixture of K 2 CO 3 and Na 2 CO 3 may be used as a flux. It is frequently expe- 
 dient to use a fusion mixture consisting of i part of NH 4 C1 and 8 parts of CaCO 3 by 
 weight. 
 
CARBON FAMILY. 243 
 
 The crucible should not be more than half-filled. Heat gently until 
 the frothing has ceased, then heat with the blast-lamp flame until the 
 decomposition is complete and the contents of the crucible are in quiet 
 fusion. When the crucible is cold it is placed upon its side in a beaker 
 containing sufficient water to cover the crucible. Hydrochloric acid is 
 now added until effervescence ceases, and no further precipitation of 
 gelatinous silicic acid takes place. Filter. Transfer the hydrogele of 
 silicic acid from the filter paper to a crucible and dehydrate it. Results ? 
 What is the substance which remains in the crucible after heating? 
 Rub some of it between the fingers. Do you detect any grit ? Equations? 
 
 Experiment IV. Fluosilicic Acid. Fluosilicates. 
 
 (a) To 5 cm. 3 of a solution of fluosilicic acid (H 2 SiF 6 ) add a small 
 quantity of a solution of potassium nitrate (KNO 3 ). Examine the 
 precipitate of potassium fluosilicate. 
 
 (b) Prove that fluosilicates are decomposed when heated with H 2 SO 4 , 
 and that silicon tetrafluoride is evolved. Recall Exp. Ill "Fluorine." 
 Record procedure. Equations? 
 
 (c) Repeat (a), using a solution of a sodium salt. Results? Equa- 
 tions ? 
 
 (d) Repeat (a), using a solution of barium chloride. Results? 
 Equations ? 
 
 (e) Record the names and formulae of the well-defined hydrides, 
 halides and acids of silicon. What is carborundum? Its formula? 
 
 Experiment V. A Test for Silicates. Silicon Tetraflouride. 
 
 (Hood.) Place 2 or 3 cm. 3 of concentrated R 2 SO 4 in a platinum 
 crucible and add a small quantity of powdered silicate or substance to 
 be tested. Warm the contents of the crucible gently; allow crucible to 
 stand until effervescence ceases. When the contents of the crucible 
 have cooled (Caution!), add i cm. 3 of H 2 F 2 . Place the cover on the 
 crucible in such a manner as to leave a small opening on one side. Hold 
 a platinum loop containing a drop of distilled water at this opening; 
 warm the crucible gently and observe whether white particles of silicic acid 
 appear in the drop of water. Evaporate the drop to dryness on the 
 crucible cover or a piece of platinum foil. A white residue which is not 
 volatilized at a high temperature, proves that silica was present in the 
 original substance. Equations? 
 
 All silicates save those of the alkali metals are insoluble in water. 
 They are decomposed by hydrofluoric acid with the formation of silicon 
 tetrafluoride. A very few of the insoluble silicates are decomposed by 
 acids other than hydrofluoric. The insoluble silicates are usually 
 converted into a soluble form by fusion with the carbonates of the alkalies. 
 
244 
 
 EXPERIMENTAL CHEMISTRY. 
 
 The physical and chemical relations of carbon and silicon may be seen 
 in the following table: 
 
 Physical Properties. 
 Atomic weight, 
 State or phase, 
 
 Color, 
 
 Specific gravity, 
 
 Specific heat, 
 Melting-point, 
 
 Chemical Properties. 
 H-derivatives, 
 Heat of formation, 
 State or phase, 
 
 Halides 
 
 Heat of formation, 
 
 State or phase, 
 
 O-derivatives, 
 Heat of formation, 
 State or phase, 
 
 S-derivatives, 
 Heat of formation, 
 State or phase, 
 
 Carbon. 
 
 12.0 
 
 Solid 
 
 (amorphous, crystalline) 
 
 Black (amorph.) 
 Black-gray (graph.) 
 
 2.2-3.5 (cryst.) 
 (graphite, diamond) 
 
 Silicon. 
 28.4 
 Solid 
 
 (amorphous, crystalline) 
 
 Greenish-brown (amorph.) 
 
 Grayish-black (cryst.) 
 
 2.4 
 (crystalline) 
 
 35 
 
 CH 4 ; C 2 H 4 , etc. 
 C, H 4 = 2i,8oocal. 
 Gas 
 
 CC1 4 ; etc. 
 
 C,C1 4 = 21,000 cal. 
 
 Liquid 
 
 CO; C0 2 
 
 C,O 2 =96,980 cal. 
 Gas 
 
 CS 2 
 
 C, S 2 = 26,100 cal. 
 
 Liquid 
 
 SiH 4 
 
 SiH 4 = 24,800 cal. 
 
 Gas 
 
 SiCl 4 ;_ 
 Liquid 
 
 cal. 
 
 SiO 2 
 
 Si,O 2 = 2 19,000 cal. 
 
 Solid 
 
 SiS 2 
 
 Si, 82 = 40,400 cal. 
 Solid. 
 
 PROBLEMS. 
 
 1. Five grams of pure graphite are completely burned in oxygen. 
 What volume of carbon dioxide is formed? What is the heat of the 
 reaction ? 
 
 2. What volume of oxygen is necessary to burn 12 grams of carbon? 
 
 3. How many grams of pure calcium carbonate will be required to 
 produce 10 liters of carbon dioxide? 
 
 4. What volume of carbon monoxide will be liberated by the action 
 of H 2 SO 4 on 10 grams of oxalic acid? 
 
 5. An analysis of the air in a lecture-room showed that 8.5 volumes 
 of carbon dioxide were present in 10,000 volumes of air. If the room 
 is 20 meters long, 17 meters wide and 7 meters high, what is the weight 
 and volume of the carbon dioxide in the room ? 
 
 6. Dumas and Stas found that 30 parts of carbon by weight com- 
 bined with So parts of oxygen, and that the carbon dioxide formed con- 
 tained its own volume of oxygen. A liter of carbon dioxide weighs 
 1.976 grams. From the foregoing data calculate the molecular weight 
 and deduce the simplest formula of carbon dioxi de. Deduce the atomic 
 weight of carbon. 
 
 7. How many liters of carbon dioxide must be passed over red-hot 
 
CARBON FAMILY. 245 
 
 charcoal to yield 100 liters of carbon monoxide? To yield 25 grams of 
 carbon monoxide? 
 
 8. A precipitate which when dry weighed 1.5 grams was formed by 
 passing carbon dioxide through lime-water. What is the weight of the 
 absorbed gas? 
 
 9. Calculate the precentage weight of CO 2 in calcium carbonate, 
 sodium bicarbonate, and magnesium carbonate. 
 
 10. How many grams of oxygen will be required for the complete 
 combustion of 5 grams of methane ? Of 5 grams of ethylene ? Of 5 
 grams of acetylene ? 
 
 11. Calculate the percentage composition of marsh gas, ethylene 
 and acetylene. 
 
 12. Water gas is prepared by passing steam over red-hot coal. The 
 simplest equation for the reaction is: 
 
 C + H 2 O H 2 + CO. 
 
 What volume of water gas at 15 C. and 730 mm. will be formed from 
 50 grams of steam? 
 
 13. Scheele found that 0.6738 of silicon tetrachloride reacted with 
 AgNO 3 to yield 2.2769 grams of silver chloride. Calculate the atomic 
 weight of silicon. 
 
 14. Calculate the percentage composition of 
 
 Garnet, Ca 3 Fe 2 (SiO 4 ) 3 
 
 Mica, KH 2 Al 3 (Si0 4 ) 3 
 
 Kaolin, H 2 Al 2 (SiO 4 ) 3 , H 2 O 
 
 Serpentine, Mg 3 Si 2 O 7 ,2H 2 O 
 
 Orthoclase (feldspar), KAlSi 3 O 8 . 
 
CHAPTER XXI. 
 
 SOME COMMON CARBON COMPOUNDS. 
 
 The simplest compounds of carbon are those which contain only hydro- 
 gen and carbon, e.g., CH 4 , C 6 H 6 , C 2 H 4 . These compounds have been 
 termed hydrocarbons and are regarded as the fundamental compounds 
 of organic chemistry. Nearly all organic compounds are either hydro- 
 carbons or hydrocarbon derivatives. Speaking comparatively, very few 
 carbon compounds are known which do not contain hydrogen. 
 
 Although hundreds of hydrocarbons are known and an almost in- 
 finite number are theoretically possible, they may be arranged in a 
 small number of comparatively simple series. These series correspond 
 somewhat to the different groups of elements. Further, the members 
 of one and the same series resemble one another much more closely than 
 do the elements of a given series of the elements. In general, the members 
 of any series bear such a close resemblance to one another, that if we 
 understand the simpler members, we are able to anticipate many of the 
 properties of the more complicated members. Attention is also directed 
 to the fact that for each hydrocarbon in a series there is a corresponding 
 class of derivatives and that the relations existing between any hydro- 
 carbon and its derivatives are very similar to those existing between any 
 other hydrocarbon of the series and its derivatives. It is obvious, then, 
 that if we know the derivatives which can be yielded by a single hydro- 
 carbon of a series, then we are able to prophesy with some degree of 
 certainty the existence of the derivatives of every other hydrocarbon of 
 the series. 
 
 The following table shows the classification of a number of the hydro- 
 carbons as regards their empirical formulae in three of the best known 
 "series." It will be observed that each hydrocarbon differs from the 
 one which precedes it by an atom of carbon and two of hydrogen. Such 
 a series is known as an homologous series and may be represented by one 
 general formula, C n H 2n+2 . 
 
 Methane Series. 
 
 Ethylene Series. 
 CH 
 
 Methane ..... C H 4 
 
 Ethane ...... C 2 H 6 
 
 Propane ...... C 3 H 8 
 
 Butane ...... C 4 H 10 
 
 Pentane 
 Hexane 
 Etc. 
 
 C 5 H 12 
 
 Ethylene 
 Propylene . 
 Butylene 
 Amylene .. 
 Hexylene.. 
 Heptylene . 
 
 246 
 
 C 2 H 4 
 
 C 3 H 6 
 
 C 4 H 8 
 
 C 5 H 10 
 
 C 6 H 12 
 
 C 7 H 14 
 
 Benzene Series. 
 C n H 2n _ 6 . 
 
 Benzene C 6 H 6 
 
 Toluene C 7 H 8 
 
 Xylene C 8 H 10 
 
 Mesitylene . . C 9 H 12 
 Pseudocumene C 9 H 12 
 
 Durene CIO^M 
 
 Cymene C 10 H 14 
 
SOME COMMON CARBON COMPOUNDS 247 
 
 Saturation. The hydrocarbons are frequently referred to as being 
 saturated or unsaturated accordingly as they behave towards chemcal 
 agents. Thus, if all the four valences of carbon are employed, the 
 hydrocarbon having then no power to combine directly with other com- 
 pounds or elements, it is said to be saturated. 
 
 CH 4 + C1 2 -> CH 3 C1 + HC1. 
 
 It is obvious that in the reaction indicated above chlorine must first 
 displace hydrogen before it can enter into combination with the com- 
 pound. On the other hand, if the hydrocarbon (any compound) can 
 unite directly with elements or compounds it is spoken of as being un- 
 saturated; for example, ethylene combines directly with chlorine to 
 form ethylene chloride 
 
 C 2 H 4 + C1 2 ->C 2 H 4 C1 2 . 
 
 Organic Radicals. The hydrocarbons may be regarded as hydrides of 
 various radicals, as methyl hydride or methane, CH 3 .H, ethyl hydride 
 or ethane, C 2 H 5 .H, etc. As a matter of fact, hydrocarbons from which 
 hydrogen has been removed give rise to hydrocarbon radicals, thus, 
 CH 3 is the organic radical methyl from methane, and C 2 H 5 is the radical 
 ethyl from ethane. These groups of units, radicals, which are found 
 to pass unaltered from compound to compound are like all other radicals, 
 incapable of existing in the jree state. Since in chemical reactions 
 these organic radicals behave toward other elements and radicals 
 in a manner similar to that in which the metals behave toward the non- 
 metals they are frequently termed positive radicals. They usually 
 differ from the inorganic radicals in that they lack the property of form- 
 ing ions. In general, organic substances may be most conveniently 
 regarded as derived from these hyrdocarbon radicals, which in most 
 cases are hydrocarbons possessing unsatisfied bonds of valency. 
 
 The following scheme, containing a partial list of the more important 
 fundamental forms of the hydrocarbon derivatives may be of service 
 in emphasizing the thought of the foregoing paragraph: 
 
 Hydrocarbons ................ R H 
 
 Alcohols ..................... R O H 
 
 Phenols ...................... R O H 
 
 Ethers ....................... R O R 
 
 Aldehydes ................... R C 
 
 Acids ....................... R C o H 
 
 Ketones ..................... R C R 
 
 Acid anhydrides .............. R C /C R 
 
 Halides . . RC1, RBr, etc. 
 
248 EXPERIMENTAL CHEMISTRY. 
 
 Esters R C \Q_ R 
 
 Amides R "" C \NH 2 
 
 Amines RNH 2 , R 2 NH 2 , etc. 
 
 Organo-Mineral Compounds ... R M 
 
 Carbohydrates C 6 H 10 O 5 , C 12 H 22 O 11 , etc. 
 
 Explanatory Note.* 
 
 Alcohols. Composed of carbon, hydrogen and oxygen and containing 
 one or more hydroxyl groups (OH), as methyl-alcohol, CH 3 OH, ethyl- 
 alcohol, C 2 H 5 OH. 
 
 Phenols. Similar to alcohols in composition, but resembling the 
 acids in many of their properties; however, they do not yield aldehydes 
 when partially oxidized as phenol, C 6 H 5 . OH. 
 
 Ethers. The oxides of the radicals; formed from the alcohols by 
 the substitution of a hydrocarbon radical for the H in the hydroxyl, 
 as methyl-ether, CH 3 . O . CH 3 ; ethyl-ether, C 2 H 5 . . C 2 H 5 . 
 
 Aldehydes. Dehydrogenated alcohols; products of the partial oxida- 
 tion of the alcohols, containing the group (COH), as formaldehyde, CH 2 O; 
 as ethyl-aldehyde (acet-aldehyde) CH 3 .COH. 
 
 Acids. Products of the further oxidation of the alcohols, containing 
 one or more carboxyl radicals, CO 2 H; as formic acid, CH 2 O 2 ; acetic acid, 
 CH 3 CO 2 H or CH 3 . CO . OH. 
 
 Ketones. (Acetone, the simplest of the ketones, is prepared by the 
 distillation of calcium acetate.) Formed from the acids by the substi- 
 tution of a hydrocarbon radical for the OH in the carboxyl; contain the 
 group CO, as acetic tetone or acetone CH 3 . CO . CH 3 . 
 
 Halides. Formed from the foregoing groups by the substitution of a 
 halogen radical for hydrogen or hydroxyl; as chloroform, CHC1 3 ; iodo- 
 jorm, CHI 3 , ethyl chloride, C 2 H 5 C1; acetyl chloride,^ CH 3 . CO . Cl. 
 
 Esters (ethereal salts). Formed from the acids by the substitution 
 of a hydrocarbon radical for the hydrogen in the carboxyl radical; as 
 ethyl acetate, CH 3 . CO . O . C 2 H 5 . 
 
 Ammonia Derivatives. Formed upon the model of ammonia, NH 3 , 
 by the substitution of a radical for hydrogen; as acetamide, NH 2 .C 2 H 3 O; 
 ethylamine, NH 2 .C 2 H 5 . 
 
 Organo-Mineral Compounds. Formed upon the type of the chlorides 
 of metals or non-metals by the substitution of hydrocarbon radicals 
 for the chlorine, as zinc ethide, Zn(C 2 H 5 ) 2 . 
 
 Compounds for which the exact structure has not been fully estab- 
 lished. They are classified according to similarity of properties, ultimate 
 composition or products of decomposition. 
 
 Carbohydrates, or compounds usually containing six, or some multiple 
 of six, atoms of carbon, together with some multiple of the group H 2 O,f 
 as starch, C 6 H 10 O 5 ; glucose, C 6 H 12 O 6 ; sugar, C 12 H 22 O n . 
 
 *Data from Bloxam. fRhamose (C6H I2 O S ) is an exception. 
 
SOME COMMON CARBON COMPOUNDS. 249 
 
 Glucosides, or compounds which yield glucose as one of their products 
 of decomposition, as salicin, C 13 H 18 O 7 . 
 
 Albuminoids and Gelatinoids, or compounds containing C, H, N and 
 O, often with small quantities of S, and sometimes of P, distinguished 
 by their tendency to putrefy when moist; albumin, fibrin and casein 
 are examples of such compounds, but they cannot at present be repre- 
 sented by satisfactory formulae. 
 
 Isomerism. One of the most interesting facts with which we early be- 
 come acquainted in the study of organic chemistry is the existence of 
 entirely different compounds having the same percentage composition. 
 Such compounds are called isomeric or briefly isomers. Isomers are 
 of two kinds: (i) If the isomers have the same molecular weight they 
 are known as metamers, as ethyl alcohol (C 2 H 5 OH) and methyl-oxide 
 ( (CH 3 ) 2 O). (2) If they have different molecular weights, as acetylene 
 (C 2 H 2 ), benzene (C 6 H 6 ) and styrene (C 8 H 8 ), they are called polymers. 
 
 Isomers usually differ in both their physical and chemical properties. 
 For example, the empirical formula C 2 H 6 O represents two substances 
 (metamers) the one a gas (methyl-oxide), the other a liquid (ethyl 
 hydroxide or alcohol). Isomeric compounds undoubtedly owe their 
 differences to different groupings or arrangements of the atoms within 
 the molecules. It frequently occurs that one of the metamers can be 
 made directly from the other. This gave rise to that variety of chemical 
 change previously named "internal rearrangement."* In 1828, Wohler 
 effected the synthesis of a well-defined organic compound, urea, CO- 
 (NH 2 ) 2 , from wholly inorganic materials. This was the first synthesis 
 by a chemist, of a typical organic compound from a substance whose 
 preparation is independent of life processes. Wohler warmed a solution 
 of ammonium cyanate, NH 4 .CNO, which is a metamer of urea, for some 
 time. On cooling the liquid, long prisms of urea were deposited. 
 
 NH 4 CNO + CO(NH 2 ) 2 . 
 
 These two compounds are entirely different as regards their chemical 
 properties. Ammonium cyanate is an ionizable salt while the latter 
 (urea) is not a salt, but unites with acids like ammonia to form salts. 
 The ammonium cyanate undoubtedly underwent an intramolecular 
 change as indicated by above formulae; at least the attempt is made to 
 explain the difference in properties of the metamers by postulating a 
 difference in the molecular structure or constitution oj the substances. 
 
 It is quite probable that a difference in structure is closely related 
 to a difference in energy-content of metamers. For example, fumaric 
 and malei'c acids have the same empirical formula, C 4 H 4 O 4 , but the 
 "heat of combustion" of the former is 1338 Kj., while that of malei'c 
 acid is 1365 Kj. This indicates that the energy-content of malei'c acid 
 is greater than that of fumaric acid. The conclusion seems justified 
 
 * One author suggests that there may be discovered 802 compounds of the 
 ormula, d 3 H 2 8. 
 
250 EXPERIMENTAL CHEMISTRY. 
 
 when it is recalled that maleiic acid is the less stable of the two compounds 
 and melts at 130, while fumaric acid melts at a temperature above 200. 
 The following structural formulae indicate a difference in the constitution 
 of the two compounds: 
 
 COOH C H CH COOH 
 
 II II 
 
 H C COOH CH COOH 
 
 Fumaric acid Male'ic acid 
 
 (C 4 H 4 4 ) (C 4 H 4 4 ) 
 
 Graphic or structural formulae are the result of an endeavor to provide 
 a "picture" of the way in which the atoms or radicals of a molecule are 
 linked together. 
 
 It is acknowledged that any success which has been or may be achieved 
 in the synthesis of organic compounds has been or will be the result of 
 the study of the structure of compounds. 
 
 The proving of the constitution of a substance represents in many 
 cases one of the most difficult problems of organic chemistry. It usually 
 involves an accurate and extended consideration of the reactions of the 
 substance under consideration. In order that the student of inorganic 
 chemistry may have some conception of the mode of procedure of a 
 chemist endeavoring to prove the constitution of a substance, the follow- 
 ing method is cited, and may be regarded as typical although it represents 
 one of the simpler cases. 
 
 Determination oj the Constitution oj Alcohol. (a) When alcohol 
 reacts with hydrochloric acid, and the mixture is distilled at a low tem- 
 perature, the products are water and a volatile liquid known as ethyl 
 chloride, C 2 H 5 C1. The interaction may be indicated by the equation, 
 C 2 H 5 OH + HC1 -> C 2 H 5 C1 + HOH. It is evident that the Cl of HC1 
 has exchanged positions with the OH of the alcohol; this leads to the 
 conclusion that alcohol is composed of at least two groups, namely the 
 ethyl radical and the hydroxyl radical, and that its rational formula is 
 C 2 H 5 .OH. 
 
 (b) Again, sodium dissolves in alcohol (C 2 H 5 OH) with an evolution 
 of hydrogen, forming a crystalline substance known as sodium ethoxide, 
 C 2 H 5 ONa. It is seen that Na has displaced one atom of H. It might 
 be supposed if an excess of Na were used, that the remaining H 5 might 
 be displaced and a compound of the composition of C 2 Na 6 O, would be 
 produced ultimately, but this is contrary to all experimental evidence. 
 Na can be substituted for only one of the six atoms of hydrogen in 
 alcohol. It is inferred then that one of the six atoms is "linked" or 
 "fitted" into the intramolecular structure in a manner different from the 
 other five. The concept is indicated by writing the formula for alcohol 
 as follows, C 2 H 5 OH. 
 
 Chemists have been unable as yet to agree on the rational formula 
 of many compounds, especially the carbohydrates, glucosides, alkaloids, 
 
SOME COMMON CARBON COMPOUNDS. 251 
 
 etc. The following structural formula for cane-sugar (C 12 H 22 O U ) has 
 been proposed: 
 
 CH 2 OH 
 I 
 CH 2 OH-CHOH.CHCHOH.CHOH-CH-O-C-CHOH.CHOH.CH-CH 2 OH). 
 
 I II ! 
 
 | o o | 
 
 Experiment I. Hydrocarbons. 
 
 (a) Recall or repeat the experiments in which methane, ethylene and 
 acetylene were prepared. Are these hydrocarbons gases, liquids or 
 solids? Write the equations for the combustion of these three com- 
 pounds. 
 
 (6) Examine small quantities of each of the following hydrocarbons 
 and give the formula (of the main components if a mixture) for each. 
 Classify them as to state: Gasoline, toluol, kerosene, naptha, benzine, 
 petroleum, paraffine, benzene, ozokerite, vaseline and gasoline. 
 
 Experiment II. (L. T.) Determination of the "Flashing-point" of 
 Kerosene. 
 
 Note. "The temperature at which oil gives off sufficient vapor to 
 form a momentary flash when a small flame is brought near its surface" 
 is known as the "flash-point." The temperature at which oil gives 
 off enough vapor to maintain a continuous flame if ignited, is ascer- 
 tained by the "fire test." The " burning- point" is about 10 C. higher 
 than the " flash-point." The safety of kerosene depends largely on the 
 absence of volatile hydrocarbons which may escape in sufficient quan- 
 tities to form an explosive mixture with air. The presence of these 
 hydrocarbons is determined by the flashing test. 
 
 (a) The simplest form of apparatus used for determining the flash- 
 point consists essentially of a small beaker in which is suspended a 
 thermometer. A small quantity of the kerosene to be tested is placed in the 
 beaker and warmed slowly till a flame which is brought near to the 
 surface at regular intervals of time, causes a momentary flash. At the 
 flashing-point the vapor ignites, and the bluish flame runs down to the 
 surface of the oil. 
 
 (b) A simple, but more accurate form of apparatus for determining 
 the flashing-point may be prepared as follows: A rubber stopper pro- 
 vided with single perforation is fitted to one end of a piece of a glass 
 cylinder 2.5 cm. in diameter, and 15 cm. long. One end of a piece 
 of ordinary glass tubing about 40 cm. long is forced through the cork 
 until it projects about i cm. beyond the inner surface of the cork; the 
 tubing is then bent off, close to the cork, twice at right angles in such 
 a manner that the long end of the tube is parallel and close to the cylinder; 
 the tubing is finally bent on a level with the open end of the cylinder at 
 right angles and away from the latter. The apparatus if made according 
 
252 EXPERIMENTAL CHEMISTRY. 
 
 to directions, has much the appearance of the old-fashioned, long-stemmed 
 pipes. The bent tube contracts to a small orifice within the cork. Air 
 is forced through this tube. The oil to be tested is poured into the 
 cylinder until the latter is filled to a point such that when air is being 
 forced through the apparatus, the surface of the foam is about 5 cm. 
 from the top of the cylinder. The apparatus is now placed in a beaker 
 of water the surface of the oil and the water should be at the same 
 level. Suspend a thermometer in the oil. Heat the contents of the 
 beaker gently and force a slow current of air through the oil. Bring a 
 small flame to the mouth of the cylinder for an instant as the temperature 
 of the kerosene rises slowly, degree by degree. The lowest temperature 
 at which the vapor ignites, as indicated by the bluish flame running down 
 to the surface of the oil, is the flash-point. 
 
 Determine the flashing-points of two or three grades of kerosene. 
 
 From what you know of the properties of gasoline, would you infer 
 that its flash-point is higher or lower than that of kerosene ? 
 
 Experiment III. Fermentation. Properties and Preparation of Al- 
 cohol. 
 
 (a) Dissolve 50 grains of grape sugar (C 6 H 12 O 6 )* in 350 cm. 3 of water, 
 and add one-half of a compressed yeast cake. Place the mixture in a 
 large flask (Fig. 45) or bottle provided with a one-hole rubber stopper; 
 
 FIG. 
 
 45- 
 
 connect the flask with a wash bottle half filled with lime-water. The 
 delivery tube of the wash bottle should be connected in series with a 
 U-tube filled with caustic potash or soda-lime, the object of which is to 
 prevent carbon dioxide in the air from acting upon the lime-water. Set 
 the apparatus aside in a moderately warm place. Fermentation usually 
 begins at once as is evidenced by the bubbling of carbon dioxide, one of 
 
 * Cane sugar (ordinary sugar) does not ferment. If boiled with acid it is inverted, 
 as indicated by the following equation: 
 
 C M H M 1X + H 2 C 6 H I2 6 + C 6 H I3 6 
 when fermentation may take place. 
 
SOME COMMON CARBON COMPOUNDS. 253 
 
 the products, through the lime-water. This action should be allowed 
 to continue for two days, then decant the liquid upon a filter. Set up a 
 condenser and a distilling flask; place the nitrate together with several 
 "boiling tubes" in the distilling flask; insert a thermometer in a hole in 
 the cork with which the distilling flask is provided, so that the bulb just 
 touches the surface of the contents of the flask; heat gently, and collect 
 that portion of the distillate which passes over between 80 and 93 C. 
 
 (b) Notice the odor of the distillate (?). Place a few drops of it on a 
 watch glass and ascertain whether it burns (?). Test its reaction with 
 litmus paper (?). 
 
 (c) Repeat (b), using (a) ethyl alcohol, (b) methyl alcohol (?). What 
 are your conclusions as to the nature of the distillate used in (b) ? What 
 is its empirical formula ? Its rational formula ? 
 
 (d) Try the solubility of iodine, camphor, rosin, shellac, etc., in sepa- 
 rate portions of alcohol (?). 
 
 C 6 H 12 O 6 2 C 2 H 5 OH + 2 CO 2 . 
 
 Ethyl alcohol, C 2 H 5 OH, is the best-known member of the class of alco- 
 hols. When pure, it is a transparent, colorless, mobile and volatile liquid 
 possessing a peculiar and agreeable odor. It has a sp. gr. of 0.797, boils 
 at 78.3 C., solidifies at 130 C., is of a neutral reaction, and burns with 
 a non-luminous flame. When mixed with water a contraction of volume 
 occurs and heat is liberated. The attraction of alcohol and water for 
 one another is so great, that strong alcohol abstracts it from tissues, 
 membranes, and other similar substances immersed in it; to this property 
 are due its coagulating action on albumin and its preservative action on 
 animal substances. Pure alcohol taken internally acts poisonously; 
 when taken in the dilute form it possesses intoxicating properties. It is 
 affirmed by various authorities that "it lowers the temperature of the 
 body from 0.5 to 2 C., although the sensation of warmth is experienced." 
 
 Alcohol which does not contain more than i per cent, of water is known 
 as pure or absolute alcohol. Diluted alcohol, made by mixing equal 
 volumes of water and alcohol, has a sp. gr. of 0.936 and is identical with 
 the proof-spirit of the United States Custom-house and Internal Revenue 
 service. 
 
 Alcohol is an excellent solvent for many organic and inorganic sub- 
 stances. These alcoholic solutions are given such names as spirits, 
 tinctures, fluid extracts, etc. The terminology depends on various phar- 
 maceutical relationships. "Spirits from the pharmaceutical point of 
 view, are simply alcoholic solutions of volatile substances." "The active 
 ingredient may be solid, liquid or gaseous." "Tinctures are alcoholic 
 solutions of medicinal substances. They differ from spirits in being 
 made from non-volatile bodies the tincture of iodine being the one 
 officinal exception to this rule." "Fluid-extracts are liquid alcoholic 
 preparations of uniform and definite strength, made by percolating 
 drugs with menstrua, and concentrating a portion of the percolate so that 
 in each case a cubic centimeter represents the medicinal virtues of one 
 
254 EXPERIMENTAL CHEMISTRY. 
 
 gram of the drug; they are mostly concentrated tinctures." "Practice of 
 Pharmacy. "Remington. 
 
 Experiment IV. Tests for Ethyl Alcohol. 
 
 (a) lodoform test. Dissolve a small flake or crystal of iodine in 2 
 or 3 cm. 3 of alcohol; add a strong solution of crystallized sodium car- 
 bonate or potassium hydroxide until the brown color of the solution dis- 
 appears; the yellow precipitate which forms is iodoform, CHI 3 . 
 
 Note. The student is reminded that other alcohols, acetone, aldehyde, 
 etc., show the same reaction. 
 
 (b) Aldehyde test. Reducing action of alcohol. To 4 or 5 cm. 3 
 of a potassium dichromate solution add i cm. 3 of sulphuric acid and 
 2 cm. 3 of alcohol; heat the solution gently and note the odor of the 
 gaseous product, acetic aldehyde (C 2 H 4 O), which is a volatile liquid 
 boiling at 20.8 C. The green color of the solution is due to the for- 
 mation of chromium sulphate, Cr 2 (SO 4 ) 3 . 
 
 (c) Ether test. Add to 2 cm. 3 of alcohol the same volume of strong 
 sulphuric acid; heat to boiling and add a little more alcohol, drop by 
 drop. The odor of ethyl ether, (C 2 H 5 ) 2 O, may be easily detected on 
 further heating. 
 
 2C 2 H 5 OH + H 2 SO 4 -* C5 >SO 4 + H 2 O. 
 
 | 5 S0 4 + C 2 H 5 OH-^ p 25 + H 2 S0 4 . 
 
 n / ^2 n 5 / 
 
 (d) Ester test. To 4 or 5 cm. 3 of a strong solution of sodium acetate, 
 NaC 2 H 3 O 2 , add a few drops of concentrated sulphuric acid and 5 or 6 
 drops of ethyl alcohol, then warm the mixture gently. Notice the 
 pleasant fruit-like odor of the ester, ethyl acetate, a volatile liquid which 
 is formed under the above conditions. The foregoing is a test for either 
 alcohol or acetic acid. 
 
 C 2 H 5 OH + C 2 H 3 O 2 H C 2 H 5 C 2 H 3 O 2 + H 2 O 
 Alcohol. Acetic acid. Ethyl acetate. 
 
 C 2 H 5 OH + NaC 2 H 3 O 2 + H 2 SO 4 C 2 H 5 C 2 H 3 O 2 + NaHSO 4 + H 2 O. 
 
 In many of their relations the alcohols are analogous to metallic 
 hydroxides. For example, CH 3 OH and C 2 H 5 OH are compounds anal- 
 ogous to NH 4 OH, KOH, NaOH, etc. Both classes of compounds react 
 with acids to form salts. Ethyl acetate, one of the products of the inter- 
 action of ethyl alcohol and acetic acid, is an organic salt analogous to 
 ammonium or sodium acetate. These organic salts are often called esters 
 or ethereal salts. The latter term is unfortunate as these salts do not 
 ionize to any extent and possess practically none of the properties of or- 
 dinary salts. 
 
SOME COMMON CARBON COMPOUNDS. 255 
 
 Experiment V. Properties of a Trihydric Alcohol (Glycerine or Gly- 
 cerol). 
 
 Place a little glycerine in a test tube. Note its color and odor (?). 
 Taste a drop of it ( ?). Rub a little of the glycerine between the fingers ( ?). 
 Test it with litmus paper (?). Ascertain its solubility in (a) water, 
 (b) alcohol, (c) ether and (d) chloroform (?). 
 
 Heat a little glycerine in a dry test tube and try to boil it. What evi- 
 dence have you that it undergoes decomposition ? Can glycerine be 
 distilled by itself ? How may it be distilled (volatilized) ? Enumerate 
 several of its uses. Write the rational formula for glycerine. 
 
 Glycerine (glycerol), C 3 H 5 (OH) 3 , is an alcohol containing three hydroxyl 
 groups. It is the trihydric or triatomic alcohol of the radical glycerl, 
 C 3 H 5 , formed by the removal of three atoms of hydrogen from the satu 
 rated hydrocarbon, propane (C 3 H 8 ), and combination of the radical 
 with three hydroxyl groups. The common animal and vegetable fats 
 and oils contain glycerine in combination with the fatty acids. These 
 compounds form a class of esters known as glycerides, which when treated 
 with alkalies undergo decomposition the fatty acids combining with the 
 metals of the alkali to form soaps, w r hilst glycerine is liberated. 
 
 Pure glycerine is a thick, colorless, odorless liquid, oily to the touch, 
 hygroscopic, neutral in reaction and rather sw r eet. It is soluble in water 
 and alcohol in all proportions, but insoluble in chloroform and ether. 
 It has a sp. gr. of 1.225, and solidifies at low temperatures forming deli- 
 quescent crystals which melt at 17 C. When heated to the boiling-point, 
 under ordinary atmospheric pressure, it undergoes decomposition; there- 
 fore, it cannot be distilled by itself. However, it can be distilled under 
 diminished pressure and is volatilized in the presence of water or when 
 hot steam is allowed to pass through it. 
 
 On account of its solvent properties glycerine is used extensively in the 
 preparation of official solutions of various organic and inorganic sub- 
 stances. These solutions have been termed glycerites. 
 
 "Nitro-glycerine" C 3 H 5 (NO 3 ) 3 is prepared by treating glycerine with 
 a mixture of concentrated sulphuric and nitric acids. The chemical 
 action is indicated by the following equation: 
 
 C 3 H 5 (OH) 3 + 3 HN0 3 - C 3 H 5 (N0 3 ) 3 + 3 H 2 O - ( - ? cal.) 
 
 It is a pale-yellow oily liquid soluble in alcohol but insoluble in water; 
 crystallizes at 20 C. in long needles and explodes very violently by 
 concussion. It can be burned in an open dish, but if heated above 250 C. 
 it explodes. One kilogram of nitro-glycerine yields after explosion about 
 725 liters of gas, measured at o C., and 760 mm. pressure. As the 
 temperature of the gas is raised to about 7000 C. by heat liberated by 
 the explosion, the volume is much larger than that suggested. The 
 explosive power of nitro-glycerine is about thirteen times as great as that 
 of gunpowder. Nitro-glycerine is the active constituent of a number 
 of explosives. Dynamite is infusorial earth impregnated with nitro- 
 
256 EXPERIMENTAL CHEMISTRY. 
 
 glycerine. While it is not readily exploded by pressure or jar, it is by 
 percussion. 
 
 Experiment VI. Properties of Phenol (Carbolic Acid). Tests. 
 
 Caution. Phenol is strongly poisonous,* and causes blisters if it 
 comes in contact with the flesh. 
 
 (a) Ask the instructor to show you a specimen of solid phenol or a 
 solution of it. Observe its empyreumatic and disagreeable odor. When 
 diluted greatly, it possesses a sweetish and afterward a caustic taste. Test 
 a dilute solution with litmus paper (?). Recall its solubility in water 
 (See Exp. XI, "Solutions"). What peculiar property of solubility in 
 water does it possess? Define "critical solution temperature." What 
 is the melting point of phenol? Indicate by rational formula that 
 phenol (phenyl hydrate) has the structure of an alcohol. 
 
 (b) To 2 cm. ^ of phenol add small quantities of a solution of potassium 
 hydroxide until the phenol is in solution. What are the products of the 
 interaction? Now add an excess of hydrochlorid acid. Results? 
 Equations ? 
 
 (c) Tests. 
 
 1. Add a few drops of a neutral solution of ferric chloride to an 
 aqueous solution of phenol. A beautiful blue color is imparted to the 
 solution. 
 
 2. Bromine water added in excess gives a yellowish-white precipitate 
 of tri-brom-phenol, C 6 H 2 Br 3 OH, which has been used medicinally under 
 the name of "bromol." 
 
 3. When phenol is heated with nitric acid it turns yellow owing, to the 
 formation of picric acid (trinitro-phenol), C 6 H 2 (NO 2 ) 3 OH. 
 
 Phenol (carbolic acid, phenyl hydrate, phenyl alcohol), C 6 H 5 OH, is 
 prepared by the distillation of coal tar. f It is one of the chief constituents 
 of the distillate obtained between I7o -i9o C. When the distillate is 
 chilled the naphthalene crystallizes out, leaving the phenol mixed with 
 various neutral oils and impurities. 
 
 The remaining liquor is treated with an alkali which dissolves the 
 phenol, forming a solution of "sodium carbolate" which separates by 
 .gravity from the undissolved neutral oils. From the solution it is pre- 
 cipitated by sulphuric acid, carbon dioxide or furnace gases the crude 
 carbolic acid separating as an oily liquid. It is purified by repeated 
 distillations and crystallizations. 
 
 Pure carbolic acid is obtained in the form of colorless, interlacing, 
 needle-shaped crystals which sometimes acquire a pinkish tint. The 
 presence of water prevents it from solidifying. It melts at 43 C. and 
 boils at 183 C. The sp.gr. is 1.065. It is very soluble in alcohol, 
 
 * Antidote. Castor oil, or olive oil, or a mixture of both, or a mixture of magnesia 
 and oil, also sodium sulphate (Glauber's salt), internally and hypodermically are used 
 as antidotes. 
 
 t Thorpe's Outlines of Industrial Chemistry. 
 
SOME COMMON CARBON COMPOUNDS. 257 
 
 ether and glycerine. At 17 C., it dissolves in 20 parts of water. It 
 also dissolves a small amount of water; therefore if phenol is shaken with 
 less than 20 parts of water at 17 C., the mixture will separate into an 
 upper layer containing about five per cent, and a lower layer containing 
 about seventy-five per cent, of phenol. At or above 68.9 C., the two are 
 miscible in all proportions. (Recall Exp. XI, "Solutions.") It is 
 very poisonous and a very powerful disinfectant or germicide. 
 
 Phenols are the hydroxyl derivatives of benzene, C H 6 . Their con- 
 stitution allies them with the alcohols, but they possess properties so 
 different from those of the alcohols of the methane and ethylene series 
 that they have been given the general but distinctive name by which they 
 are now known. They differ from the common alcohols in not yielding 
 aldehydes or acids by oxidation. Although phenol is generally called 
 carbolic acid, it has a neutral or but faintly acid reaction. It combines 
 with strong bases to form a class of salts known as the carbolates or 
 phenolates. 
 
 C H 5 OH + NaOH -> C 6 H 5 ONa + H 2 O. 
 Phenol Sodium 
 
 Carbolate 
 
 C 6 H 5 OH + HC ? H 5 p s C H 5 .C 7 H 5 O 3 + H 2 O. 
 Phenol Salicylic Salol 
 
 Acid 
 
 2C 6 H + C 8 H 4 3 -> C 20 H 14 4 + H 2 0. 
 Phenol Phthalic Phenol- 
 anhydride phthalein. 
 
 Experiment VII. Preparation and Properties of Ether. 
 
 Caution. In working with ether carefully avoid the neighborhood 
 of free flames as its vapor is easily ignited. 
 
 (a) Preparation. Recall or repeat Exp. IV (c). Equation? 
 
 (b) Pour a few cm.3 of ether into a test tube. Observe its color, 
 odor and volatility (?). Is it lighter or heavier than water? Taste a 
 drop of it (?). Allow a few drops to fall upon the back of the 
 hand (?). Pour a drop upon a watch glass. Does it evaporate 
 more or less rapidly than alcohol? Predict its boiling-point relative to 
 that of alcohol (?). Reasons for your answer? Test the ether with 
 litmus paper (?). What can you say with reference to the mutual 
 solubility of ether and water? (Recall Exp. X, "Solutions.") 
 
 To a small piece of resin or wax add several cm. 3 of ether (?). The 
 result is typical. What are your conclusions as to its solvent properties ? 
 Write the rational formula for ether (ethyl). To what class of inor- 
 ganic compounds is it analogous? 
 
 Our ordinary ether (ethyl ether) is the representative of a very large 
 class of organic compounds which are given the general name of ethers. 
 These compounds are the oxides of the organic radicals and are analogous 
 17 
 
258 EXPERIMENTAL CHEMISTRY. 
 
 to the various metallic oxides. Many ethers are products of vegetable 
 life, and have generally a characteristic and pleasant odor. Fruit 
 essences consist mainly of compound ethers (esters). That which is 
 known as the "bouquet" or " flavor" of wine and other alcoholic liquors 
 is due chiefly to ethers or compound ethers, which are formed during the 
 various stages of fermentation by the action of the acids present upon the 
 alcohol or alcohols formed. The improvement which such alcoholic 
 liquids undergo "by age" is caused by a continued chemical action be- 
 tween the substances named. Pure ethyl ether is a colorless, mobile 
 liquid which has a sp.gr. of 0.718 at 156 C. It boils at 34.6. On 
 account of its volatility, ether should be kept in strong bottles tightly 
 corked. It is easily combustible and burns with a luminous flame. 
 Ether mixes with alcohol in all proportions. Official ether contains 96 
 per cent, of ethyl ether and 4 per cent, of alcohol. It is stated by Noyes 
 (W. A.) that ether "dissolves in n.i volumes of water at 25 C., while it 
 will in turn dissolve one-fiftieth of its volume of \vater." It is a good 
 solvent for fats, alkaloids, resins and many other classes of organic com- 
 pounds. It is used extensively as an anesthetic, causing intoxication 
 and finally loss of consciousness and sensation when inhaled. It is 
 neutral in reaction. 
 
 Experiment VIII. Oxidation of an Alcohol to an Aldehyde. 
 
 (a) Acetic or ethyl aldehyde. Recall or repeat Exp. IV (b). Equa- 
 tion? 
 
 C 2 H 6 O 2H - C 2 H 4 O 
 Ethyl Acetic 
 
 alcohol aldehyde. 
 
 (b) Formic aldehyde or formaldehyde. This substance is produced 
 by passing a mixture of the vapor of methyl-alcohol and air over a heated 
 copper spiral. The experiment may be performed as follows: Pour 
 5 cm. 3 of methyl-alcohol into a test tube clamped in a vertical positon 
 then drop into the alcohol a spiral of copper wire which has been heated 
 to redness in a flame. The colorless penetrating gas w r hich is evolved 
 is formaldehyde. What is the formula for formaldehyde ? 
 
 CH 3 OH + O -* H,CHO + H 2 O. 
 
 Aldehydes are formed by the removal of hydrogen from alcohols. 
 In fact, the name aldehyde is derived from alcohol dehydrogenatum which 
 refers to its method of formation. The removal of hydrogen may be 
 accomplished by various methods, as, for example, by the oxidation of 
 alcohols, when one atom of oxygen combines with two atoms of hydrogen, 
 forming water, while an aldehyde is formed simultaneously. The 
 realtions in composition of the alcohols, aldehydes and acids are shown 
 by the following formulae: 
 
 Radicals Hydrocarbons Alcohols Ethers Aldehydes Acids 
 CH 3 CH 4 CH 3 OH (CH 3 ) 2 O C H 2 O C H 2 O 2 
 
 C 2 H 5 C 2 H 6 C 2 H 5 OH (C 2 H 5 ) 2 C 2 H 4 O C 2 H 4 O 2 
 
SOME COMMON CARBON COMPOUNDS. 259 
 
 Aldehydes, when further oxidized, are converted into acids. Many 
 of the aldehydes in consequence of their tendency to unite with oxygen 
 to form acids are strong reducing agents. Only a few of the aldehydes 
 are of practical interest, as, for examples, formaldehyde (CH 2 O or 
 H.COH), acetaldehyde (C 2 H 4 O or CH 8 .COH), the polymerized form 
 of acetaldehyde, paraldehyde ((C 2 H 4 O) 3 ) and a few others. 
 
 Formaldehyde is a gas at ordinary temperatures but may be condensed 
 to a liquid which boils at 21 C. It is readily soluble in water, and a 
 40 per cent solution has been placed on the market under the name of 
 jormalin. This same name is now given to solid paraformaldehyde, 
 (CH 2 O) 3 , which is a polymerized form of formaldehyde. It is extensively 
 used as an antiseptic. The paraformaldehyde (formalin) splits up into 
 three molecules of formaldehyde, which, escaping as a gas, is used for 
 disinfecting purposes. It does not act injuriously on the fabric or color 
 of household goods, thus possessing an advantage over chlorine and 
 sulphur dioxide. 
 
 Experiment IX. Organic Acids. Acetic Acid. 
 
 (a) To a solution of sodium acetate, NaC 2 H 3 O 2 , add a little dilute sul- 
 phuric acid. Observe the odor it may be necessary to gently warm the 
 mixture (?). Equation? Name two commercial processes by which 
 acetic acid is manufactured. 
 
 (b) Test a solution of acetic acid with litmus paper ( ?). Pour a few 
 drops of the solution upon a solution of sodium carbonate, Na^COg. 
 Results ? Equation ? What per cent, is ionized in a i N solution ? In a 
 .1 N solution? (See table). Would you infer that it is a "strong" or 
 a "weak" acid? Test its relative electrical conductivity (?). What 
 class of salts are formed by acetic acid ? 
 
 (c) Compare the odor of dilute acetic acid with vinegar (?). Try 
 the action of vinegar upon a solution of sodium carbonate ( ?). Equation ? 
 What are your conclusions as to the relation of acetic acid to vinegar? 
 
 (d) Tests for acetic acid. 
 
 1. Acetates heated with sulphuric acid evolve acetic acid which may 
 be readily detected by its odor. 
 
 2. Acetates or acetic acid heated with sulphuric acid and alcohol 
 yield the characteristic odor of ethyl acetate. (See Exp. IV. (d)). 
 
 3. Ferric chloride, when added to a solution containing acetic acid 
 or an acetate carefully neutralized, gives ferric acetate, which is a soluble 
 salt and imparts a dark red color to the solution. When this solution is 
 boiled, a brownish precipitate of basic ferric acetate separates out. 
 
 As suggested previously, many organic acids are produced by the oxi- 
 dation of alcohols. They show the characteristics mentioned of inorganic 
 acids, viz., when soluble, have an acid or sour taste, redden litmus, con- 
 tain hydrogen which is replacable by metals with the formation of salts, 
 and yield hydrogen ions. The greater number of these organic acids 
 possess these acid properties in a much less marked degree; in fact, they are 
 
260 EXPERIMENTAL CHEMISTRY. 
 
 so weak that the acid properties can often be scarcely detected. Most 
 organic acids are colorless, odorless solids; a few are liquids and scarcely 
 any are gaseous at the ordinary temperature. 
 
 Pure acetic acid, or glacial acetic acid, melts at 16.7 C. and boils at 
 120 C. It has a sp. gr. of 1.055. It causes blisters on the skin and is 
 miscible in water, alcohol and ether. It forms salts known as acetates, 
 all of which are soluble in water. (See Thorpe's " Outlines of Industrial 
 Chemistry" and Sadtler's "Industrial Organic Chemistry" for the 
 manufacture of acetic acid.) 
 
 Experiment X. Tartaric Acid. Tests. 
 
 (a) Examine some crystals of tartaric acid. Is it soluble in 
 water? What is the effect of the solution on blue litmus paper? Try 
 the effect of an aqueous solution of tartaric acid upon a solution sodium 
 bicarbonate ( ?). Equation ? 
 
 (b) Give the formula and state one well-defined use of each of the 
 following: cream of tartar, Rochelle salt, tartar emetic. 
 
 (c) Tests. 
 
 1. Tartrates are readily charred by heating them with strong sulphuric 
 acid. 
 
 2. Tartrates are decomposed (char), and evolve an odor resembling 
 that of burnt sugar when heated. 
 
 3. Silver nitrate gives with a neutral solution of a tartrate a white 
 precipitate of silver tartrate which blackens on boiling in consequence of 
 the decomposition of the salt, with separation of silver. If ammonium 
 hydroxide is added before boiling, a mirror of metallic silver will form 
 upon the glass. 
 
 Tartaric acid, H 2 C 4 H 4 O 6 , is frequently found in vegetables and fruits, 
 and especially in grapes. With the exception of the tartrates of the alka- 
 lies, all normal tartrates are either insoluble or difficultly soluble. Most 
 of the insoluble tartrates dissolve in tartaric acid forming "acid" salts. 
 
 Experiment XI. Oxalic Acid. Tests. 
 
 Caution. Oxalic acid is a poison. 
 
 (a) Repeat Exp. X (a), substituting oxalic acid for tartaric acid. 
 
 (b) Place 5 grm. of sugar in 150 cm. 3 flask; add 50 cm. 3 of nitric acid 
 (i - i); heat gently in the hood until brown fumes are given off, then re- 
 move the flame. When the brown fumes cease to escape, boil the liquid 
 down rapidly to about 10 or 15 cm. 3 ; pour the solution into a crystallizing 
 dish. Crystals of oxalic acid separate on cooling. Remove the crystals 
 and dissolve them in warm water and recrystallize them. Dissolve some 
 of these crystals in water, and add this solution to a few cm. 3 of a dilute 
 solution of potassium permanganate which has been acidulated with a few 
 drops of sulphuric acid. Is the permanganate decolorized ? Equation ? 
 
 (c) Tests. 
 
 i. When strong sulphuric acid is heated with oxalic acid or oxalates 
 carbon monoxide and carbon dioxide are evolved. 
 
SOME COMMON CARBON COMPOUNDS. 261 
 
 2. Calcium chloride gives with neutral solutions of oxalic acid a white 
 precipitate of calcium oxalate, CaC 2 O 4 , which is soluble in hydrochloric, 
 but insoluble in acetic acid. 
 
 Oxalic acid, H 2 C 2 O 4 , is a white solid, soluble in water, crystallizing 
 from the solution with two molecules of water of crystallization, H 2 C 2 O 4 , 
 2H 2 O. It occurs very widely distributed in the vegetable kingdom, as 
 in certain plants of the oxalis varieties, in the form of the acid potassium 
 salt. It may be obtained by the oxidation of many organic substances, 
 chiefly sugars, starches, etc., by nitric acid or other strong oxidizers. It 
 is prepared technically by heating sawdust or wood-shavings with a 
 mixture of caustic soda and caustic potash to about 250 C., when the 
 oxalate of these metals are formed. The mass is extracted with water, 
 and the solution evaporated to crystallization, when sodium oxalate is 
 deposited or the dissolved alkali oxalate is treated with calcium hydrox- 
 ide, insoluble calcium oxalate forming. The calcium oxalate is de- 
 composed by sulphuric acid. 
 
 Oxalic acid is probably the strongest organic acid. It is dibasic and 
 acts as a reducing agent, decolorizing solutions of the permanganates. 
 It possesses poisonous properties.* 
 
 Oxalic acid is largely used in dyeing, calico-printing and bleaching, 
 in cleaning brass and in removing iron-mould from linen. 
 
 The oxalates of the alkalies are soluble in water. All other oxalates are 
 insoluble or only sparingly soluble. 
 
 Experiment XII. Halogen Derivatives. lodoform. Chloroform. 
 
 (a) lodoform. Recall or repeat Exp. IV (a). Equation? 
 
 (b) Chloroform. The following method for its preparation is sug- 
 gested by Remsen: Mix 550 grams of bleaching powder and 1.25 liters 
 water in a 3-liter flask. Add 33 grams alcohol of sp. gr. 0.834. Heat 
 gently on a water bath until action begins. A mixture of alcohol, water 
 and chloroform will distill over. Add water and remove the chloroform 
 by means of a pipette. Add calcium chloride to the chloroform and, 
 after standing, distill on a water bath. 
 
 4C 2 H 6 O + 8 Ca(ClO) 2 2 CHC1 3 + 3 (Ca(CHO 2 ) 2 ) + 5CaCl 2 + 8 H 2 O. 
 
 (c) Note the odor of pure chloroform. Taste a drop of it (?). Test 
 its reaction with litmus paper (?). Shake a little chloroform with a 
 few cm. 3 of a solution of silver nitrate (?). Name one of the principal 
 uses of chloroform. 
 
 (d) Tests for chloroform. 
 
 1. Add a few drops of chloroform to 5 cm. 3 of Fehling's solution and 
 heat; red cuprous oxide is thrown down. 
 
 2. A strip of filter paper saturated with chloroform burns with a 
 green-mantled flame which emits vapors of hydrochloric acid. The 
 
 * Antidote. "Lime-water or calcium carbonate should be administered, but no 
 alkalies as in cases of poisoning by mineral acids, because alkali oxalates are soluble." 
 Simon's Manual of Chemistry. 
 
262 EXPERIMENTAL CHEMISTRY. 
 
 acid vapors may be rendered more visible by approaching the flame 
 with a glass rod moistened with ammonium hydroxide. 
 
 Chloroform is a heavy, colorless liquid, possessing a sp. gr. of 1.526 
 at o C.; it has a burning sweet taste; a neutral reaction, a boiling-point 
 of 61.2 C., and a melting-point of 70 C. It is very slightly soluble in 
 water, but miscible with alcohol and ether in all proportions. Pure chlo- 
 roform is not very stable. It undergoes decomposition slowly when ex- 
 posed to light and air, yielding hydrochloric acid, free chlorine, phosgene 
 (COC1 2 ) and other products. Some of the decomposition products are more 
 poisonous than chloroform, therefore when the latter is used as an anes- 
 thetic it is essential that the chloroform should be pure. It has been 
 found that the presence of a small amount (.5 to i%) of alcohol renders 
 the chloroform more stable. Commercial chloroform usually contains 
 a small quantity (the U. S. P. allows one-half to one per cent.) of al- 
 cohol for this purpose; thereby changing the sp. gr. to about 1.488. 
 Pure chloroform should not give more than a faint opalescence when 
 shaken with a silver nitrate solution. 
 
 Experiment XIII. Esters. Ethyl Acetate. 
 
 Recall or repeat Exp. IV (d). Equation? 
 
 Recall the note following Exp. IV. The esters are analogous to inor- 
 ganic salts in structure, and it is frequently convenient to refer to them as 
 "organic salts," but their resemblance to said salts ceases, however, with 
 structure. The esters react very slowly. Originally, the esters were 
 called "compound ethers"; for example, ethyl acetate was called "acetic 
 ether." The terms are still occasionally employed. 
 ,; All true jats are esters of the triatomic alcohol, glycerine, and the higher 
 members of the fatty acids. They are frequently termed glycerides ; stearin, 
 for example, being the glyceride of stearic acid. Fats are widely dis- 
 tributed in the animal and vegetable kingdoms. They are found in 
 animals generally under the skin, around the intestines and on the muscles, 
 while in plants they exist chiefly in the seeds. 
 
 Human fat, lard, mutton fat and bee/ tallow are mixtures of palmitin, 
 C 3 H 5 (C 16 H 31 2 ) 3 , stearin, C 3 H 5 (C 18 H 35 O 2 ) 3 , and olein, C 3 H 5 (C 18 H 33 O 2 ) 3 . 
 
 Butter consists essentially of the glycerides of butyric acid, caprylic acid, 
 caproic acid and capric acid which are volatile with steam, and of palmitic 
 and stearic acids, which are not volatile. Palmitin and stearin are solids, 
 but olein is a liquid at ordinary temperatures. The relative amount of 
 these three fats determines its solid or liquid condition. The chief con- 
 stituent of liquid fats is generally olein. The liquid fats have been given 
 the name jatty oils or fixed oils in contradistinction to the essential or 
 volatile oils which give the peculiar odors to plants. Among the more 
 familiar essential oils are spirits of turpentine, camphor, menthol, cedar 
 oil, oils oj cloves, lemon, mustard, peppermint, and wintergreen, attar oj 
 roses, etc. 
 
SOME COMMON CARBON COMPOUNDS. 263 
 
 The taste and color of fats are due to foreign substances which are 
 often produced by slight decomposition. Pure fats are odorless, colorless 
 and tasteless substances which stain paper permanently. They are lighter 
 than water and insoluble in it, sparingly soluble in alcohol, easily soluble 
 in ether, carbon disulphide, benzene, etc. Fats heated above 300 C. 
 decompose with the formation of various products, some of which possess 
 a very disagreeable odor. Among the products is the aldehyde, acrolein, 
 C,H 4 0. 
 
 Many of the pure fats keep without change, but since the greater num- 
 ber contain impurities, such as albuminous matter, etc., they suffer de- 
 composition,* which results in a liberation of the fatty acids. The latter 
 impart their odor and taste to the fats, causing them to become what is 
 generally termed rancid. 
 
 Drying oils are a class of fats containing unsaturated acids which 
 undergoing oxidation, render them hard. The principal vegetable dry- 
 ing oils are linseed oil, hemp oil, poppy oil and sunflower oil. 
 
 Among the vegetable semi-drying oils are corn or maize oil, cotton-seed 
 oil and castor oil. 
 
 The vegetable non-drying oils of importance are olive oil, peanut oil, 
 palm oil, cocoanut oil and almond oil. 
 
 Butterine and oleomargarine are perfect substitutes for common butter 
 when the essential attributes are considered. They are made from mix- 
 tures of animal and vegetable oils flavored with a relatively small quantity 
 of butter and colored to imitate it. Oleo oil from tallow and lard are 
 largely used. These are frequently mixed with cotton-seed oil in cold 
 weather to increase the percentage of olein. 
 
 Experiment XIV. Preparation of Soaps. Glycerine. 
 
 (a) Preparation of soda soap. Put about 10 grams of lard or tallow 
 in a 300 cm. 3 Erlenmeyer flask; dissolve 4 grams of NaOH in 100 cm. 3 
 of water, pour this solution into the flask; heat the mixture over a wire 
 gauze until it ceases to foam, then add an equal volume of a saturated 
 solution of sodium chloride. Stir constantly during the addition of the 
 salt solution. Boil for a few minutes. Allow the mixture to stand until 
 it separates into an upper solid zone or crust and a lower liquid zone 
 which may be more or less colored. The solid layer is a mixture of soda- 
 soap and unsaponified fat; the liquid contains glycerine, salt and various 
 impurities. Drain off this liquid and add 3 grams of NaOH dissolved in 
 75 cm. 3 of water; heat as previously directed until the mixture becomes 
 adhesive. Pour the mixture into an open dish. The solid product 
 should show the essential properties of hard soap. Equations? Explain 
 the precipitation ("salting out") of the soap by means of sodium chloride. 
 
 (b) Preparation of potassium soap. Dissolve 5 grams of KOH in 100 
 cm. 3 of alcohol; add 7 grams of lard; heat the mixture on a water bath until 
 it has the consistency of syrup and the odor of alcohol is no longer per- 
 ceptible. The mixture should be stirred constantly while it is being 
 
 * A kind of fermentation which is aided by oxidation. 
 
264 
 
 EXPERIMENTAL CHEMISTRY. 
 
 heated. Pour the mixture into an open dish and allow it to cool. The 
 jelly-like product is soft soap. The alcohol is used as a common solvent 
 for the fat and alkali. 
 
 C 3 H 5 (C 16 H 31 O 2 ) 3 
 Palmitin 
 (Fat) 
 
 C 3 H 5 (C 18 H 35 p 2 ) 3 
 Stearin 
 (Fat) 
 
 C 3 H 5 (C 18 H 33 O 2 ) 3 
 Olein 
 
 C 3 H 5 (C 16 H 31 O 2 ) 3 
 
 Palmitin 
 
 3 NaOH -* 3 NaC 16 H 31 2 . 
 
 Sodium palmitate 
 (Hard soap) 
 
 + 3 NaOH -. 3 NaC 18 H 36 2 
 
 Sodium stearate 
 (Hard soap) 
 
 + 3 NaOH - 3 NaC 18 H 33 O 2 
 Sodium oleate 
 
 + 3 KOH ^ 3 KC 16 H 31 2 
 
 Potassium palmitate 
 (Soft soap) 
 
 C 3 H 5 (OH) 3 
 Glycerin 
 
 C 3 H 5 (OH) 3 
 Glycerin 
 
 C 3 H 5 (OH) 3 
 Glycerin 
 
 C 3 H 5 (OH) 3 
 Glycerin 
 
 Experiment XV. Properties of Soaps. 
 
 (a) Dissolve a little of the soap prepared in Exp. XIV (a) in warm 
 water; filter. To one-half of the filtrate add hydrochloric acid and shake 
 vigorously. Observe that the fatty acids separate as solids and rise to 
 the top. Remove this floating coagulum and test its solubility in 
 alkali (?). Equations? 
 
 (b) To the other half of the soap solution add a calcium sulphate or 
 magnesium sulphate solution. Note the formation of the corresponding 
 insoluble calcium or magnesium salt of the fatty acid. Equations? 
 Explain how the efficiency of soap is diminished by using it with hard 
 water. What information do we convey by the expression sojt water ? 
 
 (c) Test yellow soaps, toilet soaps, etc., for free alkali. Tabulate your 
 results giving the name of the soap used. 
 
 As previously explained, the common fats, glycerides or esters undergo 
 decompostion when boiled with water, acids or alkalies w r ith the forma- 
 tion of the free acid or a salt of the acid, and the alcohol (glycerin). 
 This process of decomposing an ester is called saponification. The 
 metallic salts which are formed when the saponification is effected by 
 means of an alkali are called soaps. They are defined as metallic salts of 
 certain non-volatile fatty acids. Soaps intended for washing purposes 
 should contain only soluble salts of the acids; i.e., those of sodium or 
 potassium; the magnesium, calcium, lead and other heavy metal salts 
 are insoluble in water. The alkalies which are most frequently used in 
 the manufacture of soaps are sodium and potassium hydroxides. The 
 former yields a "hard soap"; the latter, a "soft soap" which is liquid 
 under ordinary conditions because of a greater solubility and a lower 
 melting-point. It is also probable that the greater deliquescence of the 
 potassium soaps contributes to their "softness." 
 
SOME COMMON CARBON COMPOUNDS. 265 
 
 Experiment XVI. Carbohydrates. Grape Sugar (Glucose). 
 
 (a) Examine some glucose (grape sugar) and tabulate its properties. 
 What is its formula? 
 
 (b) Glucose ferments. Recall or repeat Exp. Ill (a). Equation? 
 
 (c) "Fehling's Test" for glucose. Place about 10 cm. 3 of Fehling's 
 solution in a test tube; add a few drops of a solution of glucose; boil hard 
 for several minutes, then add a little more of the glucose solution. Continue 
 to boil until on removing the tube from the flame, a dark-red precipitate 
 settles to the bottom, leaving a clear colorless liquid. The red precipitate 
 is cuprous oxide. This reaction can be used both for detection and 
 quantitative determination, if other reducing substances are absent. 
 
 Experiment XVII. Carbohydrates. Cane sugar (Saccharose). Inver- 
 sion of Cane Sugar. 
 
 (a) Repeat Exp. XVI (a), substituting cane sugar or beet sugar for glu- 
 cose. 
 
 (b) Cane sugar does not ferment. Repeat Exp. XVI (b) substituting 
 cane sugar or beet sugar for glucose. Results? 
 
 (c) Inversion of cane sugar. Prepare a dilute solution of pure cane- 
 sugar (rock-candy). Divide it into two parts. Test one portion with 
 Fehling's solution as in Exp. XVI. Result? Now add a few drops 
 of concentrated hydrochloric acid to the other portion of the sugar solu- 
 tion ; heat for thirty minutes or longer, on a water bath at i oo C . ; make the 
 solution exactly neutral with a solution of sodium carbonate and test with 
 Fehling's solution. Results? Conclusion relative to the presence of 
 glucose ? Equation ? What is the function of the hydrochloric acid ? 
 
 Experiment XVIII. Carbohydrates. Preparation and Properties of 
 Cellulose. Parchment Paper. 
 
 (a) Treat a small quantity of cotton-wool or Swedish filter paper 
 successively with ether, alcohol, water, a caustic alkali, and finally, a di- 
 lute acid. Wash with water (?). Enumerate some of the properties and 
 uses of cellulose. What is its formula? 
 
 (b) Parchment paper. Immerse a strip of filter paper in a solution 
 prepared by adding 3 cm. 3 of concentrated sulphuric acid to 15 cm. 3 
 of water. Allow the paper to remain in the solution for about fifteen 
 seconds, then remove it to a large vessel of pure water and wash it thor- 
 oughly. Rinse it in a very dilute solution of ammonia and dry it in the 
 air. Result ? 
 
 Experiment XIX. Conversion of Cellulose into Glucose. 
 
 Cut a filter paper into small pieces and dissolve it in the least volume 
 of concentrated sulphuric acid; add 500 cm. 3 of water and boil for about 
 an hour. Add pulverized chalk (calcium carbonate) to remove the sul- 
 phuric acid; filter, and evaporate the filtrate to small bulk. Test for 
 
266 EXPERIMENTAL CHEMISTRY 
 
 glucose with Fehling's solution. Result ? The chemical reaction may be 
 indicated by the following equation: 
 
 (C 6 H 10 5 ) n + nH 2 -* n(C 6 H 12 6 ) 
 Cellulose Glucose 
 
 Experiment XX. L. T. Cellulose Nitrate (Gun-cotton or Smokeless 
 Powder). 
 
 (a) Mix i volume of nitric acid which has a sp. gr. of 1.4 with three 
 volumes of concentrated sulphuric acid; cool the mixture to about 12 C.; 
 add as much pure cotton, small quantities at a time, as can be sub- 
 merged in the acid mixture. Let the cotton remain in the mixture for 
 twenty-four hours, then remove it and wash it several times in large 
 volumes of water. Let it dry in the air. The white, friable mass which 
 is very unstable is gun-cotton. To a small quantity of the prepared 
 explosive touch a lighted match ( ?). 
 
 2(C 6 H 10 5 ) + 6HN0 3 -^ C 12 H 14 4 (N0 3 ) 6 + 6H 2 O* 
 Decomposition on explosion: 
 C 12 H 14 4 (N0 3 ) 6 7 C0 2 + 6CO + 5~N~+ 8H~+ 3 H 2 O + 1074 calf 
 
 (b) State briefly the composition, preparation and uses of each of the 
 following: cordite, celluloid and collodion. 
 
 Experiment XXI. Carbohydrates. Properties of Starch. 
 
 (a) Dust a glass slide with starch, then examine the starch through a 
 microscope. Make a diagram of the "field of vision." 
 
 (b) Ascertain if starch is appreciably soluble in cold water or alco- 
 hol (?). In hot water (?). 
 
 (c) Test for starch. Recall test for iodine (?). 
 
 (d) Cut a very thin section from a piece of potato by use of a micro- 
 tome or a razor; float the section in a small dish of water to which has been 
 added a few drops of an iodine solution. Remove the section from the 
 water and spread it upon a glass plate, then cover it with a "cover glass." 
 Examine the section through a microscope. Draw a diagram of what 
 you see. 
 
 Experiment XXII. Conversion of Starch into Glucose. 
 
 (a) Test a little starch paste with Fehling's solution (?). 
 
 (b) To 100 cm. 3 of starch paste add 5 cm. 3 of concentrated sulphuric 
 acid; heat the mixture on a steam bath for two hours or until iodine no 
 longer causes a blue color (which shows complete conversion of starch 
 into either dextrin or glucose), and until i cm. 3 of the solution yields no 
 precipitate when 5 cm. 3 of alcohol are added to it, dextrin being pre- 
 cipitated by alcohol. The dextrin should be converted into glucose. 
 
 * Guttmann, Industrie der Explosivstoffe. f Berthelot. 
 
SOME COMMON CARBON COMPOUNDS. 267 
 
 If the flask containing the mixture is connected with an inverted con- 
 denser, it will avoid the necessity of replacing the water which would other- 
 wise escape. Proceed from this point as in Exp. XIX. Results ? 
 
 (C 6 H 10 5 ) n + nH 2 n(C 6 H 12 6 ). 
 Starch Glucose. 
 
 Carbohydrates. These compounds are found widely disseminated in 
 the vegetable kingdom; in fact, no other organic substances are found in 
 such abundance. They are also found as products of animal life, as, 
 for example, the sugar in milk. 
 
 Most carbohydrates are white, solid substances, generally soluble 
 in water. Those belonging to the sugars have a more or less sweet taste. 
 They are of neutral reaction, and are either fermentable or can, in most 
 cases, be converted into substances capable of fermentation. Many 
 of the carbohydrates, especially glucose, are easily oxidized; i.e., they are 
 good reducing agents, as is shown in the foregoing experiments by the 
 reduction in alkaline solution of the salts of copper (Fehling's solution). 
 
 The carbohydrates are now conveniently arranged under three heads. 
 They are: 
 
 1. Monosaccharides or simple sugars. 
 
 This group contains those sugars which cannot be broken down into 
 two or more sugars. Examples of these are glucose and fructose. 
 
 2. Polysaccharides or complex sugars. 
 
 Examples are cane sugar, sugar of milk and maltose. 
 
 3. Polysaccharides, not resembling sugars. 
 Examples are cellulose and starch. 
 
 Glucose. C 6 H 12 O 6 (grape sugar, dextrose). This substance is widely 
 distributed throughout the vegetable kingdom, and is generally accompa- 
 nied by fructose or fruit sugar. It is found in relatively large quantities 
 in the juices of many fruits, such as the grape, fig, strawberry, cherry, 
 mulberry, etc., and in honey. Traces (o.i per cent, or less) of grape sugar 
 occur in the liver, in the normal blood and in the urine. In the disease 
 diabetes mellitus, the quantity secreted in the urine is greatly increased, 
 frequently rising as high as 5 to 10 per cent. Glucose is easily prepared 
 from pure cane sugar, which is easily hydrolyzed by the action of dilute 
 acids, yielding glucose and fructose as indicated by following equation: 
 
 C 12 H 22 U + H 2 -> C 6 H 12 6 + C 6 H 12 6 
 Cane sugar Glucose Fructose 
 
 Commercially, it is manufactured in large quantities by heating starch 
 with dilute sulphuric acid, which converts the starch first into dextrin and 
 maltose, and then into glucose. The solution obtained by this process 
 is treated with chalk to remove the excess of acid, and filtered. The 
 filtered solution is either evaporated to a syrupy consistency and sold as 
 "glucose," or to dryness, the solid product obtained being known as 
 commercial "grape sugar." 
 
268 EXPERIMENTAL CHEMISTRY. 
 
 Glucose is soluble in water and alcohol. Its sweetness, compared with 
 that of cane sugar, is as 3 to 5. It is fermentable and reduces Fehling's 
 solution. It is optically active, turning the plane of polarization to the 
 right. (a) D = + 52.7 for a 10 per cent, aqueous solution which has 
 been prepared for some time. Cold, freshly prepared solutions give 
 a much larger rotation. 
 
 Saccharose, C 12 H 22 O n (cane-sugar, beet-sugar). Sucrose (saccharose) 
 is formed in the juices of many plants, but it is usually associated with 
 substances which render its extraction exceedingly difficult and unprofit- 
 able. The sugar-cane, sugar-beet, sugar-maple and the date-palm are 
 the commercially important sources of sugar, the first two supplying nearly 
 all the sucrose of commerce. 
 
 Sucrose forms white, hard, crystalline granules, but may be obtained 
 by crystallization from water in large monoclinic prisms (rock-candy). 
 It is sparingly soluble in alcohol, but dissolves in 0.5 its own weight of 
 water at temperatures below 20 C., and in 0.2 parts at 100 C.; it melts 
 at 160-161; and on cooling it solidifies to a pale-yellow, amorphous, 
 glassy mass known as barley sugar, which after a long time becomes 
 opaque and crystalline; when heated to 200-210 C., it loses water and is 
 gradually converted into a brown, almost tasteless substance called 
 caramel, which is used largely for coloring liquors, soups, etc. Oxidiz- 
 ing agents act energetically upon cane sugar, converting it into oxalic 
 acid and saccharic acid; a mixture of cane sugar and potassium chlorate 
 will deflagrate with explosive violence when moistened with a drop of 
 sulphuric acid; an acid solution of potassium permanganate is decolorized 
 (reduced) by sugar. Although sucrose is a strong reducing agent, it does 
 not reduce Fehling's solution. 
 
 Pure sucrose does not ferment, but if a trace of a mineral acid is added 
 to an aqueous solution and the liquid warmed or allowed to stand, the 
 sucrose is hydrolyzed with the formation of equal quantities of the fer- 
 mentable substances, glucose and fructose. 
 
 C B H B O il + H 2 0- C 6 H I2 6 + C 6 H 12 6 . 
 
 Sucrose Glucose Fructose. 
 
 Sucrose is optically active, (a) D = + 66.51 at 20 C. for a solution 
 containing 26.004 grams of sugar in a volume of 100 cm. 3 Therefore, 
 sucrose is dextrorotatory, i.e., rotates the plane of polarized light to the 
 right, but when it has been hydrolyzed with acids as indicated by fore- 
 going equation, the resulting solution of glucose and fructose is levorota- 
 tory, i.e., the plane of polarization has been reversed or "inverted." 
 This is due to the fact that fructose* causes a greater rotation to the left 
 than glucose does to the right. The mixture of glucose and fructose is 
 called invert sugar, and the process of converting sucrose into these two 
 forms is called inversion. An instrument known as the saccharimeter 
 is used to determine the rotatory power of a known solution before and 
 
 * (a) D = 93 in 10 per cent, aqueous solution. 
 
SOME COMMON CARBON COMPOUNDS. 269 
 
 after inversion, and from this data the percentages of cane sugar and in- 
 vert sugar present in the sample can be calculated. 
 
 The ease with which sucrose is inverted indicates that the radicals of 
 glucose and fructose, which it evidently contains, are united by oxygen. 
 (See probable formula of sucrose, page 251.) 
 
 Cellulose (C 6 H 10 O 5 ) n . Cellulose is the principal constituent of cell 
 membranes and of the woody fiber of plants, and constitutes, in fact, the 
 framework of all vegetable tissues. Some parts of vegetables, such as 
 cotton, flax and hemp, consist almost wholly of cellulose. 
 
 Pure cellulose is a white, transparent mass, insoluble in all the common 
 solvents; it dissolves slowly in concentrated sulphuric acid, and if the solu- 
 tion be diluted with water and boiled, dextrine and finally glucose are 
 produced. Hence, it is possible to convert wood into sugar. 
 
 Cellulose is used extensively in the manufacture of fabrics, paper and 
 explosives. 
 
 Paper consists essentially of cellulose fibers matted or felted into a coher- 
 ent layer or sheet. Usually it is "loaded" with mineral matter to give 
 it weight and render it smooth and less porous. 
 
 Starch (C 6 H 10 O 5 ) n . Starch is found in abundance in the vegetable 
 kingdom, occurring in nearly all plants in a greater or lesser quantity. 
 It is found chiefly in the seeds, particularly of the cereals and the legumin- 
 os&, but also in the roots and stems of many plants. It is formed in the 
 leaves and green parts and then transported in soluble form to other parts 
 of the plant where it is used to build up the tissues, or is deposited as re- 
 serve material. The details of the formation of starch are not fully 
 understood, but it appears that chlorophyl acts like a "contact" substance 
 or a "catalyser" in the reaction by which starch is formed in the plant. 
 The carbon dioxide of the air is reduced by the joint action of sunlight 
 (energy), and protoplasm and the chlorophyl, the carbon being assimilated, 
 while a portion of the oxygen is set free. If the various stages in which 
 the reaction takes place are disregarded, the ultimate chemical change 
 may be represented by the following thermo-chemical equation: 
 
 6C0 2 + 5 H 2 C 6 H 10 5 + 60 2 671,000 cal. 
 or 
 
 6CO 2 + 6H 2 O -> C 6 H 12 O 6 + 6O 2 . 
 
 That substances of such a high molecular weight as starch or glucose 
 should be formed directly, as indicated by above equations, is regarded 
 by chemists as highly improbable. Baeyer* holds that formaldehyde, 
 CH 2 O, is the first reduction product of carbon dioxide in plants, and 
 that glucose from which starch may be derived is formed by a polymeri- 
 zation process: 
 
 6CO 2 + 6H 2 O 6CH 2 O + 6O 2 
 6CH 2 -* C 6 H 12 6 
 C 6 H 12 6 - C 6 H 10 5 + H 2 0. 
 
 * Berichte der deutchen chcmischen Gesellschaft, 3, 67. 
 
270 EXPERIMENTAL CHEMISTRY. 
 
 Starch as found in its natural condition possesses a granular structure, 
 the granules differing from each other so markedly that the source from 
 which they are derived can usually be determined by a microscopic exam- 
 ination. Each starch granule is made up of different layers, arranged 
 abound a nucleus which is generally at one side of the granule. The in- 
 terior substance of the granule consists of "granulose," while the trans- 
 parent, inert and insoluble covering resembles cellulose* in structure. 
 
 Ordinary starch is a white, amorphous, tasteless substance, insoluble 
 in water and alcohol, but if heated with water the granules swell and rup- 
 ture the cell walls, when the "granulose" combines with water to form 
 a jelly-like paste. If this paste is boiled with an excess of water it forms a 
 partial solution or an emulsion, which passes readily through a filter 
 paper. This solution yields a characteristic blue color with iodine, hence 
 the use of the latter as an "indicator" in many analytical operations. 
 Starch is levorotatory and is converted into dextrin or British gum by 
 heating it in the dry condition to about 200 C. When boiled with dilute 
 acids, starch is converted successively into dextrin, maltose and glucose. 
 A process which is of the utmost importance in the manufacture of alcohol 
 and spiritous liquors is the decomposition of starch at about 65 C. by 
 the ferment diastase into dextrin and maltose: 
 
 3 C 6 H 10 5 + H 2 C 6 H 10 5 + C 12 H 22 U . 
 
 Starch is an important article of food, especially when associated with 
 albuminous substances, as in ordinary flour. 
 
 In the United States starch is manufactured principally from maize 
 and wheat; in Europe, mainly from potatoes, but also from maize, wheat 
 and rice. 
 
 * "Recent investigations tend to prove that "starch cellulose" is not present as 
 such in the granule, but is formed from starch substance by the action of acids or 
 fermentation . " Thorpe. 
 
CHAPTER XXII. 
 ALUMINUM FAMILY. 
 
 Boron, B, n . o 
 (Aluminum, Al, 27.1) 
 (Earth metals,) 
 
 On inspecting the periodic table it will be seen that the fourth column 
 contains the elements, boron, aluminum and a number of rare " earth 
 metals." The two which are of particular interest to us in this work and 
 are of the most importance are boron and aluminum. The relations 
 of these two are very similar to those existing between carbon and silicon. 
 While boron is usually classed with carbon and silicon on account of its 
 close resemblance to these elements, the connection between boron and 
 aluminum is plainly revealed in the similarity of their physical and 
 chemical properties. Boron possessing the lower atomic weight, the 
 metallic character is reduced or does not appear. It is a pronounced 
 non-metal, and its oxide and hydroxide are almost exclusively acidic. 
 Its properties approximate those of the metalloids. For this reason it is 
 customary to treat boron with them, as suggested above. 
 
 Aluminum is really a perfect metal, and with its oxide and hydroxide 
 the basic properties predominate. Although the hydroxide is capable 
 of forming metallic salts with strong bases (alkalies) of the type of sodium 
 aluminate Na 3 AlO 3 , yet owing to the higher atomic weight of aluminum 
 the basic character exceeds the acidic. All the members of the aluminum 
 family are trivalent. 
 
 Boron will be treated at this time while aluminum will be taken up in 
 connection with the metals. 
 
 BORON, B. 
 At. Wt. n.o Sp. gr. 2.5-2.63. 
 
 As may be inferred from that which has preceded, boron occupies a 
 rather isolated position among the metalloids, and may best be regarded 
 as a transition element between this class and the metals. Although it 
 seems expedient to classify boron with carbon and silicon, the elements 
 which it resembles, especially in the free state, the student is reminded 
 that the group of elements most nearly related to it must be sought for 
 among the "earth metals." 
 
 The element boron has never been found in nature in the free state. 
 It occurs most abundantly in combination as boric acid and as metallic 
 
 271 
 
272 EXPERIMENTAL CHEMISTRY. 
 
 borates, among which are tincal or crude borax, Na^E^C^, boracite (mag- 
 nesium salt) and colemanite or borate spar, Ca^BgC^. 
 
 Boron is an infusible solid substance capable of existing in both the 
 amorphous and crystalline forms. The former is a greenish-black powder 
 which may be prepared by heating the oxide of boron, B 2 O 3 , with magne- 
 sium or sodium in a covered crucible. The fused mass is then boiled 
 with hydrochloric acid. The boron is separated by nitration. Crystal- 
 line boron is obtained by the fusion of aluminum with boron trioxide. 
 On account of its hardness, it is called " adamantine boron." When 
 prepared by the above process, it is not quite pure. It contains a small 
 quantity of aluminum which is probably isomorphous in this form with 
 boron. The crystals are transparent, and in their luster and hardness 
 they resemble the diamond. The amorphous variety is probably unstable 
 with respect to the crystalline. The former has a sp. gr. of 2.5; the 
 latter, 2.63. 
 
 Boron forms an unstable hydride (probably B 3 H 3 ) which resembles 
 stibine in its readiness to undergo decomposition. However, the two 
 most important compounds of boron are boron trioxide and boric acid. 
 Although the acid, H 3 BO 3 , is known in the free state, the salts of the acid 
 are not known with any dergee of certainty. The salts are formed from 
 the so-called " condensed " acids, which are intermediate products formed 
 by the dehydration of orthoboric acid. This acid looses water when 
 heated and passes into the anhydride, B 2 O 3 , which melts forming a glass- 
 like mass. This fused substance is capable of dissolving the oxides of 
 various metals, many of which impart characteristic colors to the " borax 
 bead." These phenomena serve for the detection of such metals in analy- 
 sis. The readily fusible alkali salts of boric acid are frequently employed 
 for this purpose, and for cleaning surfaces to be soldered. 
 
 Boric acid is a very weak acid, the salts of which readily undergo hy- 
 drolysis when dissolved in water. An aqueous solution of the acid con- 
 ducts electricity but slightly better than pure water. 
 
 Alcohol and boric acid interact forming the corresponding ester which 
 is very volatile. If the alcohol is set on fire the ester imparts a green color 
 to the flame. This reaction is frequently used for the detection of boric 
 acid. 
 
 Boric acid (boracic acid) is used widely as an antiseptic and a as pre- 
 servative. 
 
 Experiment I. Preparation and Properties of Boron. 
 
 Preparation and properties similar to silicon. Repeat Exp. I, " Silicon," 
 substituting boron for silicon, and boric oxide for sand or quartz. What 
 is the valency of boron ? 
 
 Experiment II. Hydrolysis of Borates. Boric Acid. Dehydration of 
 Boric Acid. Borax. 
 
 (a) Dissolve a little powdered borax (sodium tetraborate, Na 2 B 4 O 7 ) 
 
ALUMINUM FAMILY. 273 
 
 in distilled water. Test the solution with neutral litmus paper. Result? 
 Explain. Equation? 
 
 (b) Dissolve 20 grams of borax in 100 cm. 3 of hot water. Add concen- 
 trated hydrochloric acid until the solution is strongly acid, then set it 
 aside to cool. Results? Equation? 
 
 Filter and wash out the mother-liquor from the crystals with a little 
 cold water. Dry some of the boric acid crystals by pressing them be- 
 tween filter paper. Test the solubility of the crystals in cold water (?). 
 In alcohol (?). In hot water (?). Test this latter solution with litmus 
 paper ( ?). Determine its electrical conductivity relative to water. What 
 is its degree of ionization ? (See tables.) What are your conclusions as 
 to the nature of boric acid, H 3 BO 3 ? Write the structural formula for 
 H 3 B0 3 . 
 
 (c) Place about a gram of H 3 BO 3 in a small porcelain crucible; heat 
 gently at first, then strongly. Results? Equations indicating various 
 stages in the dehydration? 
 
 (d) Dissolve 5 grams of H 3 BO 3 and 6 grams of Na 2 CO 3 in 15 cm. 3 of hot 
 water (?). Set the solution aside and allow it to evaporate slowly? 
 Result ? Equation ? Enumerate some of the uses of borax. Its formula ? 
 
 Experiment III. Preparation and Use of Borax Beads. Solubility of 
 Metallic Oxides in the Bead. 
 
 (a) Make a small closed loop on the end of a platinum test wire which 
 has been previously sealed into a piece of glass tubing; heat the wire in 
 the flame of a Bunsen burner and dip it into powdered borax; heat it in 
 the flame until the white puffy mass shrinks to a small glass-like bead 
 which should be transparent. If the bead is too small to fill the loop, 
 add more borax and heat again. The bead may be removed by dipping it, 
 while hot, into cold water; the sudden cooling shatters it. 
 
 Na 2 B 4 O 7 -> 2 NaBO 2 , B 2 O 3 . 
 
 (b) Moisten a "borax bead" with a cobalt chloride or nitrate solution. 
 Heat the bead in the oxidizing flame. Notice the color of the bead when 
 hot, also when it is cold (?). The color is usually detected by looking 
 at the bead against a white object in a strong light. Is the color of the 
 bead altered by heating it in the reducing flame? 
 
 (c) Repeat (b), substituting a copper sulphate solution for the cobalt 
 solution. Tabulate results. Salts of the following may also be used: 
 nickel, iron, manganese and chromium. 
 
 The following formulae probably represent respectively the structures 
 of borax, and the compound which is formed by the union of the metallic 
 oxide and the fused borax. 
 
 18 
 
274 EXPERIMENTAL CHEMISTRY. 
 
 Na O B Na O B 
 
 / \ / \ 
 
 00 00 
 
 \ / / / 
 
 B Co B 
 
 \ \ . 
 
 O O 
 
 B 
 
 / \ 
 O O 
 
 O 
 
 o 
 
 / \ I 
 
 Co B 
 
 Na O B 
 
 O O 
 
 \ / 
 Na O B 
 
 COLORS OF BORAX BEADS. 
 
 Metal. 
 
 Chromium, 
 
 Cobalt, 
 Copper, 
 
 Iron, 
 Manganese, 
 
 Nickel, 
 Uranium, 
 
 Oxidizing Flame. 
 
 Reducing Flame. 
 
 Cold. 
 
 Oxide. 
 
 Hot. Cold. 
 
 Green 
 
 Cr 2 3 
 
 Emerald Emerald 
 
 
 
 green green 
 
 Blue 
 
 CoO 
 
 Blue Blue 
 
 Blue 
 
 Cu 2 O 
 
 Colorless Red 
 
 
 
 (opaque) 
 
 Yellow 
 
 Fe O 
 
 Bottle-green Bottle-green 
 
 Amethyst- 
 
 MnO 
 
 Colorless Colorless 
 
 red 
 
 
 (Faint rose) (Faint rose) 
 
 Red-brown 
 
 Ni 
 
 Gray Gray 
 
 
 
 (Turbid) (Turbid) 
 
 U 3 Og Green 
 
 Green 
 
 Experiment IV. Detection of Boron. Tests. 
 
 (a) Dissolve a small quantity of boric acid in 10 cm. 3 of alcohol (ethyl 
 or methyl). Pour a portion of the mixture into an inverted porcelain 
 crucible cover and set fire to it, noting the color of the flame (?). 
 
 (b) To a strong aqueous solution of borax in a test tube provided with 
 a cork carrying a glass jet, add a little sulphuric acid, then a volume of 
 alcohol equal to that of the solution. Heat the mixture and inflame the 
 vapor issuing from the jet (?). Equations? 
 
 (c) Dip the end of a platinum wire, successively, into concentrated 
 H 2 SO 4 , glycerine, and powdered borax. Hold the wire in the outer 
 edge near the bottom of a small Bunsen flame. What color is imparted 
 to the flame? 
 
 (d) Reaction with turmeric. Dip a strip of turmeric paper into a solu- 
 tion of boric acid, or a borate to which has been added a little HC1 or 
 H 2 SO 4 to liberate the boric acid. Observe the production upon the paper 
 of a characteristic red-brown stain. The color is distinguished from that 
 produced by the alkalies by the fact that when it is touched with a drop of 
 
ALUMINUM FAMILY. 
 
 275 
 
 an alkali solution, the brown color is changed to a greenish-black color, 
 but is restored to its original color and is not discharged by dilute HC1 or 
 H 2 S0 4 . 
 
 The only borates which are readily soluble in water are those of the 
 alkalies. Magnesium borate together with a few others are difficultly 
 soluble. 
 
 The similarity of boron and aluminum may be inferred from the data 
 incorporated in the appended table: 
 
 Physical Properties. 
 Atomic weight, 
 State or phase, 
 
 Color, 
 
 Specific gravity, 
 
 Specific heat, 
 
 Melting point, 
 
 Chemical Properties, 
 
 H-derivatives, 
 
 State or phase, 
 
 Halides, 
 
 Heat of formation, 
 
 State or phase, 
 
 Stability, 
 
 O-derivatives, 
 Heat of formation, 
 
 State or phase, 
 Stability, 
 
 Acids, 
 
 Stability, 
 
 Boron. 
 
 II .0 
 
 Solid 
 
 (Amorp. Cryst.) 
 Transparent (Cryst.) 
 
 2.5-2.63 
 
 Aluminum. 
 27.1 
 Solid 
 
 Silver-white 
 2.58 (Hammered) 
 
 Infusible 
 
 BH 3 
 
 Gas 
 
 BC1 3 
 
 104,000 cal. 
 
 Liquid 
 
 Unstable in water 
 
 B 2 3 
 
 317,200 cal. 
 
 Solid 
 Soluble 
 
 (H 3 B0 3 ) 
 (H B0 2 ) 
 (H 2 B 4 7 ) 
 Stable in water 
 
 657 
 
 A1C1 3 ;A1 2 C1 6 
 
 32i,9oocal. (A1 2 C1 6 ) 
 
 Solid 
 
 Stable in water 
 
 A1 2 O 3 
 388,900 cal.(A! 2 O 3 , 
 
 3H 2 0) 
 Solid 
 Insoluble 
 
 (H.A1 3 ) 
 (H Al O 2 ) 
 (H 2 A1 2 4 ) 
 Unknown in the 
 free state. 
 
CHAPTER XXIII. 
 THE METALS OR BASE-FORMING ELEMENTS. 
 
 Introductory Note. On account of certain physical and chemical 
 properties common to a large number of the elements, but wanting in a 
 greater or lesser degree in others, the elements are, as suggested pre- 
 viously, grouped into non-metals, metalloids, and metals. In our previous 
 work we have studied chiefly the non-metallic or acid-forming elements. 
 We are now ready to begin the systematic study of the metals, having 
 studied but one typical metal, namely, sodium. 
 
 The metals have already been defined as base-forming elements. They 
 are usually good conductors of heat and electricity, and are endowed 
 with a peculiar luster which is known as metallic luster. Although the 
 number of metallic elements is much larger than that of the non-metals, 
 the chemistry of the former is less diverse and more simple. 
 
 The student is again reminded that there is no sharp distinction 
 between metals and non-metals, but rather a gradual merging of one 
 class into the other, depending on whether the distinction is based upon 
 their physical or chemical properties. 
 
 Physical Properties of the Metals. 
 
 1. At ordinary temperatures metals are solids with the exception of 
 mercury. 
 
 2. The relatively high conductivity of the metals for electricity is 
 characteristic. 
 
 3. The metals when in the compact form are endowed with a metallic 
 luster. With the exception of copper and gold, the metals in compact 
 masses possess a silver-white color. Most of the metals are black when 
 in a powdered condition, magnesium and aluminum being notable 
 exceptions. 
 
 4. Many of the metals possess the property of tenacity in a very 
 considerable degree. 
 
 5. All of the metals can be obtained in the crystallized form. 
 
 6. The metals vary widely as regards specific gravity and volatility 
 and fusibility. (See tables.) Those metals which have a sp. gr. less 
 than 4 (5, and even 6, are sometimes given as the limit) are called the 
 light metals, and those with a greater sp. gr., the heavy metals. 
 
 7. Malleability is a property peculiar to most of the metals. Antimony 
 and bismuth are reduced to powder when hammered. 
 
 8. Many of the metals when in the molten condition dissolve in one 
 another forming mixtures which are sometimes spoken of as undetermined 
 compounds, but more often as alloys. These alloys frequently possess 
 properties very similar to solid solutions. They are generally more 
 
 276 
 
THE METALS OF BASE-FORMING ELEMENTS. 277 
 
 fusible than their component metals. "Wood's metal" is an alloy 
 which melts at about 66 C., and contains bismuth 4 parts, lead 2 
 parts, tin and cadmium, each i part. (See Table of Alloys.) Alloys 
 in which mercury is one of the components are known as amalgams. 
 
 Chemical Properties Characteristic of the Metals. 
 
 1. The oxides and hydroxides of the metals usually exhibit basic 
 properties. 
 
 2. The metals manifest but little tendency to form combinations 
 with hydrogen. 
 
 3. Each metal is able to function as a positive radical in a salt, and to 
 exist as a positive ion in aqueous solutions of said salt. 
 
 4. The metals show a strong attraction for the halogen group of 
 elements, forming compounds which as a rule are soluble in water but 
 are riot hydrolyzed. 
 
 There are numerous other chemical properties which are characteristic 
 of metals, and which may serve as a basis of distinction between metals 
 and non-metals, but they are not common to all the metals. The student 
 will become familiar with these various properties as the work progresses. 
 
 Occurrence oj the Metals in Nature. Certain metals are found in nature 
 free of all combination. When a metal cocurs thus it is said to be native. 
 Platinum, gold, silver, copper, bismuth, arsenic, antimony, mercury, etc., 
 are found in the native state. More often the metals are found in nature 
 in combination with one or more of the other elements, as oxides, chlorides, 
 sulphides, carbonates, sulphates, phosphates and silicates. When 
 elements or compounds, or mixtures of elements or compounds, possess- 
 ing definite physical properties, are found in nature, they are usually 
 given the name of minerals. Various rocky materials known as the 
 matrix usually accompany the minerals. The minerals, or the minerals 
 and the rocks (matrix) in which they occur, are called ores. The art of 
 extracting the metals from their ores is called metallurgy. The process 
 of extracting the metals frequently consists of heating the ore with a 
 flux which combines with the matrix and other undesirable portions, 
 giving a fusible slag. The separation of the fused slag and the molten 
 metal is easily effected. 
 
 The following outline states the forms in which the metals occur in 
 the greatest abundance, and are of chief importance, commercially: 
 
 Native. Carbonates. 
 
 Gold Barium 
 
 Platinum Calcium 
 
 Silver Copper 
 
 Iron 
 
 Chlorides (chiefly). Lead 
 
 Potassium Magnesium 
 
 Sodium Manganese 
 
 Lithium Strontium 
 
278 EXPERIMENTAL CHEMISTRY. 
 
 Silicates. Arsenic 
 
 Aluminum Cadmium 
 
 Calcium Cobalt 
 
 Lithium Lead 
 
 Magnesium Mercury 
 
 Potassium Molybdenum 
 
 Sodium Nickel 
 
 Silver 
 
 Oxides. Zinc 
 Aluminum 
 
 Copper Sulphates. 
 
 Chromium Barium 
 
 Iron Calcium 
 
 Manganese Lead 
 
 Tin Strontium 
 Zinc 
 
 Phosphates. 
 
 Sulphides (chiefly). Calcium 
 Antimony 
 
 Classification oj the Metals. There has been much discussion as to 
 the best classification of the metals. The question cannot be answered 
 categorically and it is well to remember that for purposes of study, the 
 metals may be grouped variously, according to the selection of those 
 properties which are made the basis of comparison. Valence alone 
 may serve for classification, and in such case the arrangement will be 
 very similar to that of the periodic system. Again, the scheme of classi- 
 fication may be based on the analytical behavior of the metals. This 
 latter arrangement will bring together in many cases those metals be- 
 longing to one group of the periodic system, but in a few instances the 
 elements of a periodic group are separated. It is evident, then, that 
 any one scheme of classification must be one-sided as emphasizing certain 
 similarities more strongly than others. 
 
 The following grouping has been suggested by Ostwald: 
 
 Non-metals. Metals. 
 
 Hydrogen and the halogens. Alkali metals. j i^ght 
 
 The oxygen group. Alkaline earth metals. > 
 
 The nitrogen group. Earth metals. 
 
 The carbon group. The iron group. j fj eavy 
 
 The argon group. The copper group. 
 
 Other metals. 
 
 In the work which follows the metals will be considered in the order 
 in which they occur in families of the periodic grouping. This arrange- 
 ment affords a general view of the metals as brought together in the 
 natural groups of elements: 
 
THE METALS OF BASE-FORMING ELEMENTS. 279 
 
 Elements oj Group I: 
 
 Family M. (alkali metals). Lithium, sodium potassium, rubid- 
 ium, caesium and the radical ammonium, NH 4 , which is known as 
 the "hypothetical" metal. 
 
 Family m. Copper, silver and gold. 
 Elements oj Group II: 
 
 Family M (alkaline earth metals). Beryllium (glucinum), Mag- 
 nesium, calcium, strontium, barium and radium. 
 
 Family m. Zinc, cadmium and mercury. 
 Elements oj Group III: 
 
 Family M (earth metals). Scandium, yttrium, lanthanum, 
 ytterbium. 
 
 Family m. Boron, aluminum, gallium, indium, thallium. 
 Elements oj Group IV: 
 
 Family M. Titanium, zirconium, cerium, thorium. 
 
 Family m. Carbon, silicon, germanium, tin, lead. 
 Elements oj Group V: 
 
 Family M. Vanadium, columbium, and tantalum. 
 
 Family m. See Nitrogen Family. 
 Elements oj Group VI: 
 
 Family M. Chromium, molybdenum, tungsten and uranium. 
 
 Family m. See Oxygen Family. 
 Elements of Group VII: 
 
 Family M. Manganese. 
 
 Family m. See Chlorine Family. 
 Transitional Elements : 
 
 Of the First Long Period. Iron, cobalt and nickel. 
 
 Of the Second Long Period. Ruthenium, rhodium and pallad- 
 ium. 
 
 Of the Fourth Long Period. Osmium, iridium and platinum. 
 
CHAPTER XXIV. 
 ALKALI METALS. 
 
 Lithium, Li, 7.03 
 (Ammonium, NH 4 ,) 
 Sodium, Na, 23.05 
 Potassium, K, 39.15 
 Rubidium, Rb, 85.5 
 Caesium, Cs, 132.9 
 
 The metals of this family, together with their corresponding com- 
 pounds, bear a very close resemblance to one another. This family 
 includes those metals which are chemically the most active. The activity 
 increases with increased atomic weight. The metals tarnish very quickly 
 when exposed to the air owing to rapid oxidation. They decompose 
 
 FIG. 46. 
 
 water violently, liberating hydrogen. The oxides and hydroxides have 
 strongly basic properties. Their salts with the "active" acids do not 
 undergo hydrolysis when in aqueous solution. They are univalent, and 
 are never found in anions. 
 
 The compounds of ammonium will be discussed with those of potas- 
 sium to which they show a marked similarity. 
 
 280 
 
ALKALI METALS. 281 
 
 LITHIUM, Li. 
 
 At. Wt. 7.03 Sp. Gr. 0.59. 
 Experiment I. Flame Color of Lithium Compounds. 
 
 Clean a platinum wire by alternately dipping it into hydrochloric acid 
 and heating in the Bunsen flame until it gives apparently no color to the 
 flame, then dip the wire into a dilute solution of lithium chloride and hold 
 it in the Bunsen flame. Notice the color of the flame and examine it 
 with the spectroscope (Figs. 46 and 47) (Instructions) (?). Make a 
 
 FIG. 47. Single Prism Spectroscope. 
 
 diagram of the "scale" showing the relative positions of the sodium or D 
 lines (yellow) and the lithium lines. Nearly all flames in the laboratory 
 show the yellow lines owing to the presence of sodium. 
 
 All the common salts are readily soluble in water except the carbonate, 
 phosphate and oxide which are soluble with difficulty. 
 
 SODIUM, Na. 
 
 At. Wt. 23.05 Sp. Gr. 0.97. 
 Experiment I. Properties of Sodium. 
 
 (a) Recall from previous work, or repeat Exp. I., "Sodium." Read 
 the accompanying note. (See " Acids, Bases and Salts.") 
 
 1. Tabulate its most obvious physical properties. 
 
 2. Is sodium harder or softer than most of the metals with which 
 you are familiar? 
 
282 EXPERIMENTAL CHEMISTRY. 
 
 3. Is it lighter or heavier than water? 
 
 4. Does sodium tarnish when exposed to the air ? Explain. 
 
 5. Indicate by an equation the interaction of water and metallic 
 sodium. 
 
 6. Enumerate the chemical and physical properties which indicate 
 that sodium is a metal. 
 
 7. What is the flame color of sodium compounds? 
 
 (b) Examine at least ten different salts of sodium (end shelf), noting 
 the obvious physical properties. Give the name, formula and color of 
 each salt. Determine the solubility of these compounds. Tabulate 
 the foregoing data. 
 
 Experiment II. Preparation and Properties of Sodium Hydroxide. 
 
 (a) Dissolve about 25 grams of soda ash (Na 2 CO 3 ) in 100 cm. 3 to 125 
 cm. ^ of water in a large beaker, and heat on a wire gauze to boiling. To 
 10 grams of fresh quicklime which has been powdered, add sufficient 
 water to form a thin paste "milk of lime." Use heat if necessary to 
 start the slaking action. Add the milk of lime gradually, and with constant 
 stirring to the boiling solution. (Why?) Continue to boil for several 
 minutes ( ?), and stir constantly. Remove the flame, and allow the pre- 
 cipitate to settle. With a pipette or a small tube remove a little of the 
 liquid to a test tube, and test for the presence of a carbonate; if found 
 to be present add more milk of lime and boil; but if absent, decant 
 the liquid into a bottle; cork, and allow it to stand until the suspended 
 solid matter settles. Test a portion of the precipitate for the presence 
 of a carbonate ( ?). Siphon off (or filter rapidly) the clear supernatant 
 liquid by means of a glass siphon filled with water. A portion of the 
 solution may be evaporated to dryness and the solid substance; the 
 remaining portion should be preserved in a tightly-stoppered bottle (?). 
 
 (b) Take a little of the prepared solution between the fingers (?). 
 Test it with litmus paper ( ?). Try the flame test ( ?). Add an excess 
 of the solution to small quantities of very dilute solutions of each of the 
 following substances in separate test tubes: ferrous sulphate (?); ferric 
 chloride ( ?) ; zinc sulphate ( ?) ; mercuric chloride ( ?) ; copper sulphate ( ?). 
 Boil the contents of the tube (?). Describe the color and structure* 
 of each of the precipitates and state the affect of boiling same. 
 
 (c) Repeat (b), using the reagent labeled "NaOH Solution" (?). 
 What is your inference as to the identity of the solution prepared in (a) ? 
 By what other names is sodium hydroxide known ? 
 
 (d) Write equations representing the reactions involved in (a) and (b). 
 
 (e) Examine the sodium hydroxide (solid) as found on the side-shelf. 
 (See instructor.) Expose a small piece to the air on a watch glass 
 for several hours (?). Is it efflorescent or deliquescent? 
 
 (/) How can sodium hydroxide be converted into sodium chloride? 
 Equation ? 
 
 * Gelatinous, flocculent, curdy, pulverulent, granular or crystalline. 
 
ALKALI METALS. 283 
 
 Experiment III. Purification of Sodium Chloride. 
 
 Note. Common salt usually contains such impurities as calcium 
 and magnesium chlorides, sodium sulphate, calcium sulphate, etc. 
 The presence of the chlorides, magnesium and calcium, causes the salt 
 to become moist, especially in damp weather. 
 
 Prepare a saturated solution of sodium chloride by grinding 50 grams 
 of salt with 150 cm. 3 of water in a mortar. Filter the solution into a 
 beaker and conduct hydrogen chloride into it. It is suggested that this 
 gas be prepared by placing a small handful of salt in a generating-flask, 
 covering it with concentrated hydrochloric acid, and allowing concen- 
 trated sulphuric acid to drop slowly from a dropping-funnel into the 
 mixture. The evolved gas is passed into the saturated salt solution by 
 means of an inverted funnel or thistle tube which is attached to the 
 delivery tube of the generator. The mouth of the funnel should dip 
 just below the surface of the solution. Pure salt separates out as the 
 operation continues, but the impurities remain in solution. (Why?) 
 Explain the precipitation of the salt by hydrogen chloride. Indicate 
 the action by means of "ionic" equations. When the precipitation 
 has continued for sometime and considerable salt has separated, remove 
 the generator, and allow the salt to settle, then decant the fluid and wash 
 the salt with 15 or 20 cm. 3 of cold water, and again decant. After this 
 process of washing has been repeated several times, the last traces of 
 water may be removed with a filter pump or by throwing the salt upon 
 filter paper and pressing the salt with a spatula. Dry the salt by warming 
 it in a porcelain dish while stirring constantly with a glass rod. 
 
 Describe the appearance of the purified salt. Preserve it in a small 
 bottle for future use. 
 
 Experiment IV. Preparation of Sodium Carbonate (Solvay Process). 
 
 To 80 cm. 3 of water add 20 cm. 3 of ammonium hydroxide. Using 
 the foregoing mixture as a solvent, prepare a saturated solution of am- 
 monium carbonate, then place the solution in a corked bottle and 
 saturate it with sodium chloride. This can be done only by prolonged 
 agitation with the powdered salt. Allow the suspended matter to settle, 
 then decant the clear liquid into a bottle fitted with a cork and a delivery 
 tube which reaches to the bottom; connect the delivery tube with the 
 laboratory supply of carbon dioxide and saturate the solution with the 
 gas. This latter operation may require nearly an hour. Finally filter 
 off the precipitated matter, and dry the salt by pressing it between filter 
 papers. Dissolve a little of the salt in a few cm. 3 of water. Dip a clean 
 platinum wire into the solution and hold it in the colorless Bunsen 
 flame (?). Add hydrochloric acid to the solution and test the evolved 
 gas for carbon dioxide ( ?). What are your conclusions as to the identity 
 of the salt ? 
 
 Experiment V. Purification of Sodium Carbonate by Crystallization. 
 
 Using 100 grams of sodium carbonate, prepare a saturated solution of 
 
284 EXPERIMENTAL CHEMISTRY. 
 
 the salt at 32 C. ; filter at the same temperature and collect the filtrate 
 in a large porcelain dish; cool the solution to o C. by packing the dish 
 in broken ice; heavy layers of white crystals soon form in the porcelain 
 dish. Decant the supernatant liquid when the crystals are no longer 
 deposited, then dry the crystals as thoroughly as possible by pressing 
 them between layers of filter paper. Redissolve these crystals in water 
 at 33 C. and repeat the foregoing process. The crystals are subjected 
 to this treatment a third time, after which they are pressed and dried. 
 The molecular formula of the salt is now Na 2 CO 3 . ioH 2 O. Place the 
 crystals between watch glasses, and allow them to remain exposed to the 
 air for several days. They effloresce, and in a comparatively short 
 time are transformed into white powdery Na 2 CO 3 . Preserve the salt in a 
 clean glass-stoppered bottle. The relative solubility of sodium carbonate 
 at various temperatures is indicated in following table: 
 
 At 32, 100 parts of water dissolve 1140 parts. 
 At 30, 100 parts of water dissolve 273 parts. 
 At 20, 100 parts of water dissolve 92 parts. 
 As o, 100 parts of water dissolve 21 parts. 
 
 Experiment VI. Dehydration of Hydrous Sodium Carbonate. 
 
 Place about a gram of hydrated sodium carbonate in a test tube clamped 
 so that its mouth is inclined slightly downward; heat tube to a dull red- 
 ness, and note the change in the salt. What is evolved? Dissolve a 
 small quantity of the anhydrous salt in a few cm. 3 of water, and add a 
 few drops of hydrochloric acid to the solution ( ?). Equations ? 
 
 Experiment VII. Action of Acids upon Sodium Carbonate. 
 
 Try the action of each of the following acids upon separate portions 
 of a solution of sodium carbonate; hydrochloric acid (?); nitric acid(?); 
 sulphuric acid (?); acetic acid (?). Represent the reactions by ionic 
 equations. 
 
 Experiment VIII. Hydrolysis of Sodium Carbonate. 
 
 Test a solution of purified sodium carbonate with litmus paper (?). 
 Explain the reaction by the assistance of equations. 
 
 Experiment IX. Preparation and Decomposition of Sodium Hydro- 
 gen Carbonate. 
 
 (a) Preparation. Make a saturated solution of sodium carbonate by 
 warming (avoid boiling) it with water in a beaker; add an excess of the 
 salt; allow the solution to cool, filter, and pass carbon dioxide into the 
 filtrate until a precipitate forms ( ?). Indicate the reaction by an equa- 
 tion. Is sodium hydrogen carbonate more or less soluble than sodium 
 carbonate ? Test the solution with litmus paper and explain its reaction. 
 
 (b) Decomposition. Place about i gram of sodium hydrogen car- 
 bonate (bicarbonate of soda) in a test tube, add 10 cm. 3 of water and 
 heat gently. Test the evolved gas for carbon dioxide (?). Equations? 
 
ALKALI METALS. 285 
 
 Experiment X. Preparation of Sodium Hydrogen Sulphate. 
 Recall or repeat Exp. X. "Neutralization Salts." 
 Experiment XI. Reactions of Sodium Salts. 
 
 Note. Use the purified sodium chloride for the following reactions. 
 
 (a) What color is imparted to the flame by sodium salts? Examine 
 the sodium flame by means of a spectroscope. Locate its spectrum 
 on the scale (?). 
 
 (b) To a few drops of a strong aqueous solution of the chloride on 
 a watch glass add five or six drops of a platinic chloride, PtCl 4 , solution. 
 Evaporate carefully to small volume. Allow the mixture to cool, and 
 determine the solubility of the red-colored monoclinic crystals (?), (a) 
 in water, and (b) in alcohol (?). 
 
 (c) Sodium cobaltic nitrite, Co(NO 2 ) 3 .3NaNO 2 *, produces no pre- 
 cipitate in acetic acid solutions of sodium salts. 
 
 (d) Sodium salts are but slightly volatilized when heated in a porcelain 
 dish over a Bunsen flame. 
 
 All sodium salts are soluble. 
 
 Experiment XII. (Quant.) Alkalimetry and Acidimetry. 
 
 (a) Determination of the strength of a NaOH solution by titration. 
 Calculate the weight of pure (99.95 per cent, purity) oxalic acid, C 2 H 2 O 4 - 
 . 2H 2 O, required to make 500 cm. 3 of a . 5N oxalic acid solution. Powder 
 and expose on a watch-glass for 20 min. a quantity of the acid slightly 
 in excess of the calculated amount, then weigh out on glazed paper the 
 exact amount. Transfer the oxalic acid to a 500 cm. 3 graduated flask 
 which has been previously calibrated, and fill the flask to the mark with 
 distilled water. Fill a burette with the solution. Measure carefully 
 into a small flask from another burette 10 cm. 3 of the sodium hydroxide 
 solution prepared in Exp. II, dilutef with twice its volume of water, and 
 add a few drops of a phenolphthalein solution. Now cautiously run acid 
 into the flask until the red color just disappears. Record the exact 
 amount of acid used, and calculate (a) the weight of sodium hydroxide 
 per liter, (b) the normality of the solution. 
 
 (b) If the sodium hydroxide solution is above normal, calculate the 
 volume of water and solution necessary for the preparation of 100 
 cm. 3 of a normal solution. If the solution is less than normal calculate 
 the amounts required for 250 cm. 3 of a half-normal solution of the alkali. 
 
 (c) Using the amounts calculated in (b) , prepare either a normal or 
 half-normal solution of the alkali. 
 
 (d) Using the solution prepared in (c), determine the normality of the 
 * Sodium cobaltic nitrite may be prepared as follows: To 33 grams of sodium ni- 
 trite add 100 cm.3 of water, make slightly acid with acetic acid, then add 3.5 grams of 
 cobalt nitrate. Allow the solution to stand for several hours, and filter if not clear. 
 The solution decomposes slowly. 
 
 f A concentrated solution may decompose the indicator, phenolphthalein. 
 
286 EXPERIMENTAL CHEMISTRY. 
 
 hydrochloric acid on the reagent shelf. Use methyl-orange (instructions) 
 for an indicator. 
 
 How can the amount of acetic acid in vinegar be determined ? 
 
 (e) Define alkalimetry. Acidimetry? 
 
 POTASSIUM, K. 
 
 At. Wt. 39.15 Sp. Gr. 0.87. 
 Experiment I. Properties of Potassium. 
 
 (a) Observe the same caution in the manipulation of potassium as 
 suggested in the use of sodium. Place a piece of potassium upon a 
 dry paper and cut off a piece the size of a small pea. Observe the color 
 and luster of the freshly cut surface (?). Note the effect of air (?). 
 Why is it kept under kerosene ? Is the potassium hard ? Half fill a 
 tall beaker with water which reacts neutrally. Scrape the coating from 
 the piece of metal, and throw it (metal) upon the water contained in 
 beaker which should be covered with a glass plate. Does the metal 
 float ? Is there any evidence of chemical action ? Compare its interac- 
 tion with water with that of sodium under similar conditions ( ?). 
 
 (b) Wrap a piece of potassium in a little paper; place it in a wire 
 gauze basket; hold it under water and collect the gas which escapes in a 
 test tube by displacement of water. Apply a lighted match to the 
 mouth of the tube. Results? Identify the gas (?). 
 
 (c) Test the water in the beaker with red litmus paper (?). Take a 
 little of the water between the fingers (?). Dip a clean platinum wire 
 into the water and hold it in the colorless Bunsen flame. Is the flame 
 colored? What is the color? 
 
 (d) Pour 3 cm. 3 of the shelf reagent labeled potassium hydroxide 
 into a test tube and dilute with an equal volume of water. Repeat (c), 
 using the foregoing solution. Results? Inference? Indicate the 
 interaction of potassium and water by an equation. What is the valency 
 of potassium? 
 
 (e) Place a drop of bromine in a porcelain dish and cautiously drop 
 a fragment of potassium into it. Results ? Equations ? 
 
 (/) Put a small piece of potassium in a test tube with a little iodine. 
 Result ? Heat the tube gently. Results ? Equation ? 
 
 (g) Enumerate the chemical and physical properties which indicate 
 that potassium is a metal. 
 
 Experiment II. Potassium Salts. 
 
 Examine at least ten different salts of potassium, noting the obvious 
 physical properties. Give the name, formula, color and solubility of 
 each salt. Tabulate the foregoing data. 
 
 Experiment III. Preparation and Properties of Potassium Hydroxide. 
 
 Proceed as in Exp. II, "Sodium." What are your conclusions as to 
 
ALKALI METALS. 287 
 
 the similarity of the chemical properties of sodium hydroxide and 
 potassium hydroxide? 
 
 Experiment IV. Action of Acids upon Potassium Hydroxide. 
 
 (a) Preparation of potassium sulphate. Dissolve 5 grams of potassium 
 hydroxide in 20 cm. 3 of water. Add two drops of a phenolphthalein 
 solution, then neutralize the solution with dilute sulphuric acid. Evapo- 
 rate until crystallization begins, then allow the solution to cool and 
 note the formation of crystals. Equation ? 
 
 (b) Preparation of potassium hydrogen sulphate. Prepare a solution 
 of potassium hydroxide as in (a), neutralize with a measured quantity of 
 dilute sulphuric acid, then add a volume of the dilute acid equal to that 
 used in neutralizing the alkali. Evaporate to small bulk and cool. If 
 a crystalline mass is not obtained, evaporate further and cool. Repeat 
 if necessary. Equation ? 
 
 (c) Repeat (V), using nitric instead of sulphuric acid. Equation? 
 
 (d) Repeat (a) , substituting hydrochloric for sulphuric acid. Equation ? 
 
 Experiment V. Hydrolysis of Potassium Carbonate. 
 
 Repeat Exp. VIII, "Sodium," using potassium carbonate instead of 
 sodium carbonate. 
 
 Experiment VI. Preparation of Potassium Nitrate by Metathesis. 
 Relative Solubilities of Sodium Chloride and Potassium Nitrate. 
 
 Dissolve 20 grams of sodium nitrate in 40 cm. 3 of hot water. To the 
 boiling solution add 18 grams of potassium chloride. Stir with a glass 
 rod until all the salt is in solution. Evaporate to one-half this volume. 
 What separates from the boiling mixture? Rapidly filter the mixture 
 by decanting the hot clear liquid from the deposit of salt upon a filter. 
 Set the filtrate aside to be examined later. Now throw the deposited 
 salt upon the filter and press out the mother liquor with a spatula. 
 Recrystallize this salt. Compare these crystals with any that may have 
 formed in the filtrate from the original solution. Identify the two kinds 
 of crystals (?). 
 
 The process may be understood more clearly by inspecting the solu- 
 bilities of sodium chloride and potassium nitrate, the products of the 
 interaction of the factors. The following table gives the number of 
 grams dissolved by 100 cm. 3 of water: 
 
 AtoC. 10 C. 100 C. 
 
 Sodium chloride, 35.6 grm. 35.63 grm. 39.9 grm. 
 
 Potassium nitrate, 13.3 grm. 25.00 grm. 247.0 grm. 
 
 Experiment VII. Deflagration. Gunpowder. 
 
 (a) Pulverize separately and mix intimately 3 grams of potassium 
 nitrate and i gram of sulphur. Throw the mixture into a red-hot iron 
 crucible or ignite it on an iron plate. Results? Dissolve the residue 
 in water. Devise a method for testing for potassium sulphate (?). 
 
288 EXPERIMENTAL CHEMISTRY. 
 
 (b) Repeat (a), using powdered charcoal instead of sulphur. Test the 
 residue for potassium carbonate (?). 
 
 (c) Powder separately and mix thoroughly on paper 6 grams of potassium 
 nitrate, i gram of sulphur and i gram of charcoal. Place the mixture on 
 an iron plate or brick in the hood and ignite it. What is gunpowder? 
 Define deflagration. 
 
 Experiment VIII. Preparation of Potassium Cyanide and Potassium 
 Thiocyanate. 
 
 SeeExp. XVI, "Carbon." 
 
 Experiment IX. Reactions of Potassium Salts. 
 
 Note. Use chemically pure potassium chloride for the following 
 reactions: 
 
 (a) What color is imparted to the flame by potassium salts ? Examine 
 the potassium flame by means of a spectroscope. Locate its spectrum 
 on the scale with reference to the sodium or D lines. Make a diagram. 
 
 (b) To a few drops of a strong aqueous solution of the chloride on 
 a watch glass add three or four drops of a solution of platinic chloride, 
 PtCl 4 . Results? Equation? Determine the solubility* of the yellow 
 crystalline precipitate (a) in water and (b) in alcohol. 
 
 (c) Sodium cobaltic nitrite, when added to concentrated solutions of 
 potassium salts which contain acetic acid but no free inorganic acid, 
 gives a yellow pulverulent precipitate. Equation? 
 
 (d) Potassium salts are but slightly volatilized when heated in a porce- 
 lain dish over a Bunsen flame. 
 
 Most potassium salts are soluble in water. 
 
 AMMONIUM, NH 4 . 
 
 Experiment I. Preparation and Properties of Ammonium Hydroxide. 
 
 RecallExp. II, "Nitrogen." Repeat Exp. II, (b), "Sodium." Recall 
 or determine the relative electrical conductivities of normal solutions of 
 potassium hydroxide, sodium hydroxide and ammonium hydroxide. 
 (See table, Degree of lonization.) What ions do they possess in common ? 
 What are your conclusions as to the relative activities of these three bases ? 
 Would you class ammonium hydroxide among the active bases? Why? 
 Can ammonium be obtained in a free state? In what state or condition 
 does it exist ? What name is frequently applied to it ? Why ? 
 
 Experiment II. Preparation of Ammonium Salts. 
 Recall Exp. IV, "Nitrogen." 
 
 * Potassium chloro-platinate (potassium platinic chloride) is soluble in no parts 
 of water at 10 C.; nearly insoluble in alcohol, and quite insoluble in a mixture of 
 alcohol and ether. 
 
ALKALI METALS. 289 
 
 Experiment III. Hydrolysis of Ammonium Chloride. 
 
 Test a solution of ammonium chloride with litmus paper. Explain 
 result. Equation ? What is sal ammoniac ? 
 
 Experiment IV. (Optional.) Preparation of the Unstable Ammonium 
 Amalgam. 
 
 Place about 50 cm. 3 of a saturated solution of ammonium chloride in 
 a dish and warm gently. Remove the flame and add a small quantity 
 of sodium amalgam (see Appendix) to the solution while warm. Results ? 
 Hold a piece of red litmus paper over the dish (?). Equation ? 
 
 Hg x . Na + NH 4 C1 Hg x . NH 4 + NaCL 
 2 Hg x .NH 4 - 2 Hg x + 2NH S + H 2 . 
 
 Experiment V. Properties of Ammonium Salts. 
 
 (a) Dissociation. In a small dry evaporating dish heat successively 
 small quantities of the chloride, the nitrate and the sulphate. Note the 
 odor of the fumes. What is the common result ? Explain the formation 
 of the white cloud which appears above the heated salt. Is the reaction 
 reversible? (Recall Exp. Ill, "Chemical Equilibrium.") Equation? 
 
 (b) Decomposition. Heat a small quantity of ammonium phosphate 
 in a hard glass tube (?). Dissolve the residue in water and test with 
 litmus paper. Results? Conclusions? Equation? Are all ammon- 
 ium salts completely volatilized by simple heating? Is ammonia liber- 
 ated from all salts of ammonium when they are heated ? (Recall Exp. I, 
 "Nitrogen and the Atmosphere," and Exp. XI, (a), "Nitrogen," before 
 answering this last question.) 
 
 (c) Purification of ammonium chloride. Plug the stem of a large glass 
 funnel with clean cotton-wool and invert it over the dish in which a 
 mixture of sand and ammonium chloride is being heated. Continue to 
 heat the salt until a heavy white layer of it accumulates oh the inside of 
 the funnel. Scrape the salt on to a piece of glazed paper and examine 
 it. How does it compare in purity with the original mixture ? Bottle 
 the salt and save for use in Exp. VI. 
 
 Exp. VI. Reactions of Ammonium Salts. 
 
 (a) Do ammonium salts impart color to the flame ? 
 
 (b) Place a small quantity of an ammonium salt solution in a beaker 
 or a flask and add a few cm. 3 of a solution of sodium hydroxide or 
 potassium hydroxide. Cover the beaker with a watch glass on the 
 under side of which is placed a moistened piece of turmeric paper or 
 red litmus paper. Heat the beaker gently. Result. Equation ? 
 
 (c) To a few drops of a strong aqueous solution of the chloride on a 
 watch glass add three or four drops of a solution of platinic chloride. 
 Result? Equation? Determine the solubility of the yellow crystalline 
 precipitate (a) in water and (b) in alcohol. 
 
 This yellow precipitate may be distinguished from the similar potas- 
 19 
 
2QO EXPERIMENTAL CHEMISTRY. 
 
 sium salt by the fact that when heated strongly it leaves a residue of 
 spongy platinum. 
 
 (d) Sodium cobaltic nitrite precipitates from acetic solutions of am- 
 monium salts a yellow pulverulent precipitate. Equation ? 
 
 (e) Nessler's solution (a solution of potassium mercuric iodide in 
 potash) is used to detect small traces of ammonia in natural waters. 
 It gives a brown precipitate, or imparts a color according to the quantity 
 of ammonia present. 
 
 2(HgI 2 ,2KI) + 3 KOH + NH 3 NHg" 2 I ,H 2 O + 7 KI + 2 H 2 O. 
 
 (/) Ammonium salts undergo dissociation or are decomposed when 
 heated in a porcelain dish. 
 
 All ammonium salts are soluble in water. Reagents can form pre- 
 cipitations only in concentrated solutions. 
 
 Experiment VII. Detection of the Alkali Metals in a Mixture. 
 
 How can the salts of sodium and potassium be distinguished from 
 ammonium salts without the use of the flame and the spectroscope? 
 How could you distinguish between the salts of sodium and potassium 
 without the aid of the flame test or spectroscope? Devise a system 
 of tests which will provide for the detection of the three alkali metals 
 in a solution containing their respective salts. Apply the system to a 
 " known" solution of their salts. Ask the assistant to give you an 
 "unknown" solution. Test for the presence of the alkali metals. Re- 
 port (?). Make a complete record of all \vork. 
 
 RUBIDIUM, Rb. 
 
 At. Wt. 85.5 Sp. Gr. 1.52. 
 
 See text-book and lecture notes. 
 
 OESIUM, CS. 
 
 At. Wt. 132.9 Sp. Gr. 1.88 
 See text-book and lecture notes. 
 
 Lithium, rubidium, and caesium are usually placed in the category 
 of rare metals. Lithium salts cost about 10 s. per pound while rubidium 
 and caesium salts can be obtained for 64 per pound. 
 
 PROBLEMS. 
 
 1. Calculate the percentage of sodium sulphate in Glauber's salt, 
 Na 2 SO 4 .ioH 2 O. 
 
 2. What weight of sodium carbonate is theoretically required to 
 prepare 100 grams of sodium hydroxide? 
 
 Na 2 CO 3 + Ca (OH) a - 2 NaOH + CaCO 3 . 
 
ALKALI METALS. 291 
 
 3. How much sal soda (washing soda), Na 2 CO 3 . ioH 2 O, can be 
 made from i kg. of sodium chloride? The equations are: 
 
 2NaCl + H 2 SO 4 Na 2 SO 4 + 2 HC1. 
 
 CaCO 3 + C Na 2 CO 3 + CaS + 4 CO. 
 
 4. How much potassium carbonate is necessary to prepare 100 grams 
 of potassium hydroxide ? 
 
 5. How many grams of sodium hydroxide are necessary to liberate 
 the ammonia from 50 grams of ammonium nitrate ? From 50 grams 
 of ammonium sulphate ? 
 
CHAPTER XXV. 
 
 Copper, Cu. 63 . 6 
 Silver, Ag. 107.93 
 Gold, AU. 197.2 
 
 The elements of this family occur free in nature, and on this account 
 are among those metals which were earliest known. They are used 
 very extensively for ornamental purposes, and almost universally for 
 coinage. 
 
 These metals present many additional properties which are in marked 
 contrast to those of the alkali elements. They are among the least 
 active metals. They do not interact with oxygen or water at ordinary 
 temperature. Silver and gold, together with the platinum family, are 
 known as the " noble metals" because of their greater resistance to the 
 chemical influences of air and water and their easy recovery from com- 
 bination by means of heat. The oxides and hydroxides, with the ex- 
 ception of silver oxide, Ag 2 O, possess weakly basic properties. The hal- 
 ides of these metals, save those of silver, are hydrolyzed. 
 
 Upon the basis of the periodic classification these three elements fall 
 into the same family, but in many of their chemical relations they are very 
 dissimilar. For example, silver is uniformly univalent, forming one series 
 of salts; copper is both univalent and bivalent, forming two series; gold 
 is univalent and trivalent, and likewise forms two series. Univalent 
 copper and gold bear a marked resemblance to mercury, but bivalent 
 copper resembles manganese (Mn"), zinc (Zn"), iron (Fe") and 
 nickel (Ni"), and trivalent gold bears a resemblance to aluminum (Al"') 
 and iron (Fe'"). On the other hand, silver chloride, AgCl, and cuprous 
 chloride, Cu 2 Cl 2 , are both soluble in ammonia but insoluble in water. 
 
 It is quite obvious that that family is not homogeneous, yet in many 
 of their physical attributes these metals show a regular gradation in their 
 properties. As regards ductility and malleability, silver is intermediate 
 between copper and gold, the latter possessing these properties in the 
 maximum degree. With respect to tenacity, silver is also intermediate, 
 gold being the least tenacious of the three metals. These metals are the 
 three best conductors of electricity. The electrical conductivity of silver 
 exceeds that of all other metals. 
 
 COPPER, CU. 
 
 At. Wt. 63.6 Sp. Gr. 8.95. 
 Experiment I. Properties of Copper. 
 
 (a) Procure a piece of copper wire, and clean the surface by scraping 
 it with the blade of an old pen-knife or by scouring it with sand. Ob- 
 
 292 
 
COPPER. 293 
 
 serve and state its most obvious physical properties. Is copper a good 
 conductor of electricity? Of heat? 
 
 (b) Oxidation of copper. Repeat or recall Exp. XVII, (e), " Carbon." 
 Equation ? What change is slowly undergone by copper when exposed 
 to the air? Upon what experimental evidence is your answer based? 
 
 (c) Solution tension. Place 5 cm. 3 of a solution of any compound of 
 mercury in a test tube, then dip a clean copper wire into the solution. 
 After a short time, remove the wire and rub it carefully with a soft cloth. 
 Result ? Has any of the copper passed into solution ? Explain. Equa- 
 tion? Compare the action with that which takes place when zinc acts 
 upon sulphuric acid. 
 
 Insert an iron nail or wire and a strip of zinc into separate test tubes 
 half-filled with a copper sulphate solution. After a short time remove 
 the metals and examine them. Result? Has the color of the copper 
 sulphate solution undergone any change ? Account for results ? Equa- 
 tions ? 
 
 (d) Solubility in acids. Using small portions of copper filings in sepa- 
 rate test tubes, determine the solubility of the metal in hydrochloric acid. 
 In nitric acid diluted with half its volume of water ? In strong sulphuric 
 acid ? In aqua regia ? Use heat, if necessary, in each case ? Equations ? 
 What color is common to the various solutions? Recall Exp. IV, (b}, (c). 
 "The Modern Theory of Solution." How does the electrolytic dissocia- 
 tion theory account for the similarity in color of the solutions enumerated ? 
 
 (e) Make a list of the alloys of which copper is one of the components. 
 (See Appendix.) 
 
 Experiment II. Reduction of Cupric Oxide. 
 
 (a) Recall Exp. VIII, "Hydrogen." Equation? 
 
 (b) Thoroughly mix on paper about 2 grams of cupric oxide and o.i 
 gram of powdered charcoal; place the mixture in a test tube and heat. 
 Test the evolved gas for oxygen and carbon dioxide, and examine the 
 residue in the tube. Results ? Equation ? 
 
 Experiment III. Preparation and Properties of Cuprous Oxide. 
 
 To 10 cm. 3 of Fehling's solution in a test tube add 4 or 5 cm. 3 of a 
 dilute solution of glucose, or a solution of cane sugar which has been 
 boiled with a few drops of hydrochloric acid for 10 minutes. (?). Boil 
 the mixture* vigorously for several minutes. The red precipitate is cu- 
 prous oxide. Explain its formation. 
 
 Divide the precipitate into three parts. Expose one portion to the 
 action of the air (?). To the second portion add strong nitric acid and 
 boil (?). Equation? Add ammonium hydroxide to the third portion, 
 and shake the solution in a test tube. Result? 
 
 * The solution must be strongly alkaline. 
 
294 EXPERIMENTAL CHEMISTRY. 
 
 Experiment IV. Preparation and Properties of Cupric Hydroxide. 
 Cupric Oxide. 
 
 To 3 cm. 3 of a solution of copper sulphate add an excess of a solution 
 of sodium hydroxide. Results? Write the reaction. Boil one-half 
 the mixture and account for the change in the color of the precipitate. 
 Equation ? To the other half of the mixture add ammonium hydroxide. 
 Result? Equation? Write the formula of the complex ion which 
 imparts the deep blue color to this solution. 
 
 Experiment V. Ammonio-cupric Compounds, Cuprammonium. 
 Ammonio-cuprous Compounds, Cuprosammoniuni. 
 
 (a) Add ammonium hydroxide, at first in small amounts, then in ex- 
 cess, to a solution of cupric sulphate. Results? Give the formula for 
 ammonio-cupric sulphate. The copper now forms a part of what ion ? 
 Equations? Would you infer that the concentration of the cupric ion 
 in the latter solution is greater or less than in the copper sulphate solution ? 
 
 (b) Place 10 or 15 cm. 3 of a cuprammonium sulphate solution in a 
 large test tube, add an excess of copper filings or turnings, cork air-tight, 
 and set aside until the solution has become colorless. Account for the 
 change in color of the solution by use of an equation. What is the for- 
 mula of the colorless complex ion? Now remove the cork from the 
 tube and shake the solution with air. Results ? Equation ? 
 
 Experiment VI. Preparation and Degree of Hydration of Cupric 
 Chloride. 
 
 (a) Recall Exp. I, (d). Indicate the interaction of copper and hydro- 
 chloric acid by an equation. 
 
 (b) Put 2 to 3 grams of cupric oxide in a test tube or small flask, add a 
 little hydrochloric acid which has been diluted with an equal volume 
 of water, boil, and continue to add acid until all of the copper oxide is 
 dissolved. What is the color of the solution? Equation? 
 
 Place three-fourths of the solution in an evaporating dish, evaporate, 
 and crystallize. State color and formula of crystals. To one-half of 
 the remaining solution add water slowly and observe the color (?). To 
 the other portion add concentrated hydrochloric acid and observe the 
 color. Account for any change of color. Equations? What is the 
 color of anhydrous cupric chloride ? 
 
 (c) Write the formulae for the following salts of copper: Verdigris, 
 blue vitriol, paris green, cupric nitrate, Scheele's green, copper acetate. 
 
 Experiment VII. Formation and Behavior of Cuprous Chloride. 
 
 (a) To 3 grams of cupric chloride dissolved in 15 cm. 3 of water in a 
 small flask, add 5 cm. 3 of concentrated hydrochloric acid and about 
 7 grams of copper filings or turnings; an excess of copper must be used. 
 Boil the mixture until it is colorless, or a drop of the brownish-black 
 liquid gives a white precipitate when added to a test tube full of water. 
 
COPPER. 295 
 
 Pour the liquid into a large volume of water in a beaker. Wash the 
 white precipitate by decantation, and keep it under water until used. 
 What change have the cupric ions undergone ? Equation ? 
 
 (b) Expose a little of the cuprous chloride covered with water, in 
 a test tube, to the sunlight. Result ? The reaction is said to be repre- 
 sented by the following equation: 
 
 l CuCl 2 + Cu. 
 
 (c) Cuprous chloride dissolves in concentrated hydrochloric acid 
 giving the two colorless complex acids, HCuCl 2 and H 2 CuCl 3 . Equa- 
 tions? What is the complex anion formed as the result of the above 
 reaction? Do complex ions tend to dissociate? Show by ionic equa- 
 tions the various equilibria involved in a solution* of cuprous chloride 
 in hydrochloric acid. 
 
 (d) To a hydrochloric solution of cuprous chloride add a little concen- 
 trated nitric acid. The resulting color of the solution shows what ion 
 to be present? Explain. 
 
 Define oxidation in terms of the electrolytic dissociation theory. 
 
 (e) Cuprous chloride is soluble in ammonium hydroxide giving a color- 
 less solution which quickly turns to a deep-blue owing to the oxidizing 
 action of the air. Equations? 
 
 (/) Precipitation of hydrated cuprous oxide. Add an excess of a sodium 
 hydroxide solution to a hydrochloric acid solution of cuprous chloride. 
 Result? Equation? Divide the mixture into two portions. Boil one 
 portion. Result? Equation? Shake the other portion with air in a 
 test tube. Result ? Equation ? 
 
 Experiment VIII. Precipitation of Cuprous Iodide. 
 
 Add a small quantity of a solution of potassium iodide to 10 cm. 3 
 of a very dilute solution of cupric sulphate. Observe the color of the 
 solution and the precipitate ( ?). Filter and wash the precipitate. Divide 
 the nitrate into two parts. Add one part to a dilute cold starch emul- 
 sion (?). Shake the other part with a little chloroform or ether (?). 
 Equations ? 
 
 Experiment IX. Dehydration of Hydrated Cupric Sulphate. 
 
 See Exp. VIII, (a), " Water." Powder a small crystal of cupric sulphate, 
 and dehydrate it by heating it gently in a test tube clamped in a nearly 
 . horizontal position. Results? Equation? What is the effect of a little 
 water upon the anhydrous salt? 
 
 Experiment X. Decomposition of Cupric Nitrate. Preparation of 
 Pure Cupric Oxide. 
 
 Dissolve 5 grams of cupric nitrate in 25 cm. 3 of distilled water, filter off 
 any insoluble particles, evaporate the filtrate to dryness in the hood, 
 
 * The solution of cuprous chloride in hydrochloric acid or ammonium hydroxide is 
 a valuable reagent used in gas analysis for the absorption of carbon monoxide. 
 
296 EXPERIMENTAL CHEMISTRY. 
 
 and heat the dry residue in a porcelain crucible until it is of a uniform 
 black color. Equation? 
 
 Experiment XI. Precipitation of Basic Cupric Carbonate. 
 
 To a solution of cupric sulphate add a slight excess of a sodium car- 
 bonate solution ( ?). Explain the participation of water in this reaction. 
 Boil the mixture with an excess of the precipitant (?). Equation? What 
 is malachite? Its formula? 
 
 2 CuS0 4 + 2 Na 2 C0 3 + H 2 < Cu 2 (OH) 2 CO 3 
 
 2 Cu 2 (OH) 2 CO 3 + Cu 4 O 3 (OH) 2 + CO 2 + H 2 O. 
 
 Experiment XII. Precipitation of Cupric Ferrocyanide. 
 
 Add a solution of a potassium ferrocyanide to a very dilute solution of 
 cupric sulphate. Result ? Equation ? 
 
 Experiment XIII. Preparation of a Double Salt. Potassium-Cupric 
 Sulphate. 
 
 Using 5 grams of powdered cupric sulphate, prepare a saturated solution 
 at 70. Calculate the weight of powdered potassium sulphate which must 
 be taken to give the same fraction of its molecular weight. Dissolve this 
 weight of salt in its own weight of water at 70, and add a few drops of 
 sulphuric acid. Mix the two solutions, taking care to secure complete 
 solution of both salts before mixing. Pour the mixture into a beaker or 
 crystallizing dish and set it aside to crystallize. Examine the whitish- 
 blue crystals and compare them with those of cupric sulphate ( ?). Give 
 the molecular formula of the crystals. Dry the crystals between layers 
 of filter paper and preserve them in a stoppered bottle for future use. 
 
 Dissolve a part of the crystals in water. Is the cupric ion present in 
 the solution? By what experimental evidence can you support your 
 answer ? 
 
 Do "compound salts" yield complex ions? 
 
 Experiment XIV. (Quant). Equivalent of Copper. 
 
 Dissolve an accurately weighed amount (about 2 grams) of cupric sul- 
 phate in distilled water in a beaker. Record the weight. Procure 
 a small rod or bar of pure zinc, file and scour it until its surface is 
 smooth and clean; wash and dry it, then weigh accurately. Record the 
 weight. Place the zinc in the solution and allow it to remain there until 
 the latter becomes colorless. Take the zinc out of the solution, carefully 
 remove the brown deposit ( ?) by directing a small stream of water against 
 it; dry and weigh it. Record the weight. What weight of zinc has gone 
 into solution ? Assuming that all of the copper has been displaced from 
 the solution, calculate from the formula the weight of copper deposited 
 upon the zinc. This weight of copper will be the equivalent of the 
 weight of zinc in solution. 
 
 The atomic weight of zinc is 65.4. Calculate the atomic weight of 
 copper from above data. 
 
COPPER. 297 
 
 Experiment XV. (Quant.) Determination of the Percentage of 
 Copper in Cupric Sulphate Crystals by Electrolysis. 
 
 Dissolve an accurately weighed amount of pure cupric sulphate 
 in water in a weighed platinum crucible or dish. Use 10 cm. 3 of water 
 for each .1 gram of salt. Add 10 drops of concentrated nitric acid to 
 make the solution thoroughly acid, and to prevent the copper from 
 depositing in a loose, spongy form instead of a hard, smooth plate. Cover 
 the crucible with a perforated watch glass, and suspend a flat spiral 
 of platinum in the solution. Place the crucible on an electrolysis stand. 
 (Instructions.) Connect the stand with the source of electricity (a bat- 
 tery of four to eight gravity cells may be used) in such a manner that the 
 platinum spiral becomes the anode and the crucible the cathode. The 
 density (ND 100 ) of the current should not be greater than i ampere, nor 
 should the electrode tension exceed 2.5 volts. As the current continues 
 to act, observe the changes which occur in the crucible (?). To deter- 
 mine when the deposition of the copper is complete, a drop of the liquid 
 is taken out from time to time after about five hours, and tested on a 
 porcelain surface with a drop of potassium ferrocyanide which has been 
 acidified with hydrochloric acid. When the electrolysis is finished, the 
 crucible is removed, carefully and repeatedly washed with water and then 
 with small quantities of alcohol, and dried in an air bath at 90 C. De- 
 siccate and weigh. What is the weight of the copper deposited ? Calcu- 
 late the percentage of copper by weight in cupric sulphate crystals. Com- 
 pare with that calculated from the molecular formula. 
 
 Explain electro-plating in terms of the electrolytic dissociation theory. 
 
 Experiment XVI. Analytical Reactions of Cupric Salts. 
 
 (a) Dip a platinum loop into a solution of any cupric salt and hold it 
 in the Bunsen flame. The compound imparts to the flame a brilliant 
 blue color which instantly changes to green when the supply of acid is 
 exhausted. 
 
 (b) Mix a small quantity of any cupric salt with sodium carbonate and 
 heat strongly upon charcoal in the reducing flame (?). Equation? 
 
 (c) The borax bead test. Copper compounds give what color to it 
 (a) when heated in the oxidizing flame; (b) in the reducing flame? 
 Observe in each case whether the color of the bead while hot is identical 
 with that when the bead is cold. 
 
 (d) Prepare a dilute solution of pure cupric sulphate and test the reac- 
 tion of the solution with litmus paper (?). Explain the result. 
 
 Do the salts of active metals undergo hydrolysis? What is your in- 
 ference as to the chemical activity of copper as a metallic element ? 
 
 (e) Pass hydrogen sulphide through 5 cm. 3 of a dilute solution of cupric 
 sulphate (?). Filter and divide the precipitate into three parts. 
 
 Determine the solubility of cupric sulphide in sulphuric acid (?). 
 
 Write the reaction for the precipitation of cupric sulphide from cupric 
 sulphate with hydrogen sulphide. Is the reaction reversible? Upon 
 what experimental evidence do you base your answer? 
 
298 EXPERIMENTAL CHEMISTRY. 
 
 Ascertain if cupric sulphide is soluble in (a) nitric acid ( ?), (b) yellow 
 ammonium sulphide (?). Equations? 
 
 (/) To a dilute solution of cupric sulphate add a potassium cyanide 
 solution (poison) drop by drop. Observe the precipitation of the un- 
 stable cupric cyanide, which readily decomposes into cyanogen and 
 cuprous cyanide. Equations? 
 
 Is cuprous cyanide soluble in an excess of potassium cyanide ? Does 
 the resulting solution show the color of the cupric ion ? Is the cupric ion 
 present ? If not, of what ion does the copper now form a part ? What 
 general name is given to this class of ions? Equation? 
 
 Pass hydrogen sulphide through a solution of cupric sulphate which 
 has been decolorized by the addition of an excess of potassium cyan- 
 ide (?). Do results indicate the absence or presence of the cupric ion? 
 Explain. 
 
 (g) The deep-blue color imparted to solutions containing copper salts 
 by the addition of an excess of ammonium hydroxide, and the deposition 
 of copper upon strips of iron or zinc when immersed in solutions of salts 
 of the former are tests frequently used. Equations? 
 
 Experiment XVII. Experimental Study of Compound Salts and Com- 
 plex Salts. 
 
 Using the crystals preserved from Exp. XIII, prepare a solution of 
 potassium-cupric sulphate. Is this a double salt or a complex salt? 
 Is the cupric ion present in the solution ? Give a reason for your answer. 
 Pass hydrogen sulphide through the solution ( ?). Do the results support 
 your previous conclusion as to the presence of the cupric ion ? Explain. 
 
 Recall or repeat Exp. XVI, (/). 
 
 Explain in terms of the ionic hypothesis the difference between " double 
 salts" and "complex salts." 
 
 The scheme of ionization for "double salts" is identical with that of 
 "complex salts." 
 
 Double salt, K 2 SO 4 .CuSO 4 .6H 2 O: 
 
 K 2 Cu(S0 4 ) 2 < 2 K- + Cu(S0 4 )" 2 (I) 
 
 Cu(SO 4 )" 2 <=> CuSO 4 + SO" 4 (II) 
 
 Cu(SO 4 ) < Cu" + SO" 4 (III) 
 
 Complex salt, K 4 Fe(CN) 6 + 4K/ + Fe(CN)"" (F) 
 
 Fe(CN)"" 6 <z> Fe(CN) 2 + 4 CN' (IF) 
 
 Fe(CN) 2 <=Fe" + 2CN' (IIF) 
 
 The ionization corresponding to equations (II) and (III) are practi- 
 cally complete for double salts at infinite dilution, and for complex salts 
 under similar conditions the reactions corresponding to equations (IF) 
 and (IIF) take place only to a very limited extent; in fact, so small, that 
 our ordinary chemical tests fail to indicate the presence of the Fe" ion. 
 
SILVER. 299 
 
 The results of refined measurements indicate that simple ions exist 
 to a very small extent in dilute solutions of complex salts and that com- 
 plex ions are found in solutions of double salts, except at great dilution. 
 
 It is obvious that the distinction between "double" and "complex" 
 salts is one of degree and not of kind. 
 
 SILVER, Ag. 
 
 At. Wt. 107.93 Sp. Gr. 10.5. 
 Experiment I. Properties of Silver. Preparation of Silver Nitrate. 
 
 (a) Examine a small piece of silver and make a record of the most 
 obvious physical properties of the metal. 
 
 Does silver undergo oxidation when exposed to the air? Upon what 
 observation do you base your answer? 
 
 (b) Dissolve a small piece of silver in nitric acid diluted with an equal 
 volume of water ( ?). Evaporate the solution to small bulk and crystallize. 
 Dry the crystals and place them in a bottle. Use these crystals for pre- 
 paring solutions of silver nitrate. Indicate the interaction of nitric acid 
 and silver by means of an equation. 
 
 Experiment II. Electrolytic Deposition of Silver. "Silver Tree." 
 
 Dissolve about 15 grams of silver nitrate in 40 cm. 3 of water. Place the 
 solution in a small glass beaker or crystallizing dish. Introduce platinum 
 electrodes into the solution so that they will be diametrically opposite to 
 one another. Join these electrodes to the battery wires. Allow the ac- 
 tion to continue for 30 minutes or longer. A beautiful deposit of silver 
 will be made on the cathode. This deposit which spreads from the 
 cathode toward the anode is called the "silver tree" because of its strong 
 resemblance to vegetable growth. 
 
 Experiment III. Replacement of Silver in its Compounds by Copper. 
 
 Place a strip of iron, zinc, tin, lead or copper in a solution of silver 
 nitrate contained in a test tube. Result ? Ionic equation ? Examine 
 the table of "Solution Tensions." Name the above metals in the order 
 of increasing solution tension. Does silver ordinarily displace hydrogen 
 from aqueous solutions of the acids? 
 
 Experiment IV. Precipitation of Silver Oxide. 
 
 To a dilute solution of silver nitrate add a sodium hydroxide solu- 
 tion ? Result ? Equations ? Is the precipitate soluble in excess of the 
 precipitant ? 
 
 Experiment V. Preparation and Behavoir of the Halides of Silver. 
 
 (a) Silver chloride. To 5 cm. 3 of a dilute solution of silver nitrate add 
 a slight excess of dilute hydrochloric acid (?). Equation? Filter and 
 wash the precipitate with water. 
 
3OO EXPERIMENTAL CHEMISTRY. 
 
 Put a small part of the precipitate upon a watch glass and expose it to 
 direct sunlight (?). Equation? 
 
 To a second portion of the precipitate add ammonium hydroxide ( ?). 
 Equation ? What complex ion is formed ? Now add dilute nitric acid 
 drop by drop in excess (?). Equation? 
 
 Treat the third portion of the precipitate as in the preceding case, using 
 an ammonium carbonate solution instead of Ammonium hydroxide. 
 Equations ? 
 
 (b) Silver bromide. Repeat (a), substituting a potassium bromide solu- 
 tion for the hydrochloric acid. 
 
 (c) Silver iodide. Repeat (a), using a solution of potassium iodide in 
 place of the hydrochloric acid. 
 
 What extensive practical use is made of the change produced in silver 
 salts by light ? Do the halides of silver show any gradation as to color ? 
 If so state it? 
 
 Experiment VI. Preparation of Pure Silver from an Alloy. 
 
 Dissolve a copper-silver alloy in dilute nitric acid in a beaker or cas- 
 serole. Use heat if necessary. Does the solution possess any property 
 which indicates the presence of copper? Explain. Evaporate just to 
 dryness, and take up the residue with warm water. Filter, add a few 
 drops of nitric acid to the filtrate, and while stirring add dilute hydro- 
 chloric acid in slight excess. Continue to stir the mixture until the major 
 portion of the precipitate has collected together. Filter and wash with 
 water containing a few drops of nitric acid. Place the precipitate in a 
 beaker, put several pieces of granulated zinc on it, and add sufficient dilute 
 sulphuric acid to cover contents. The mixture should be stirred from 
 time to time. Equation? 
 
 After several hours, pour off the acid, remove any unchanged zinc, 
 wash the precipitate with water by decantation, and filter. Dry the 
 brown powder between sheets of filter paper, then fuse it on a stick of 
 charcoal by directing the flame of a blast-lamp downward upon it (?). 
 Scour the metal bead with sea sand ( ?). 
 
 What is sterling silver? What is the composition of standard silver 
 coin? 
 
 Experiment VII. Precipitation of Silver Sulphide. 
 
 (a) Pass hydrogen sulphide through 2 or 3 cm. 3 of a dilute solution of 
 silver nitrate (?). Equation? Is this reaction reversible theoretically? 
 To a portion of the product add dilute nitric acid. What are your con- 
 clusions now as to the reversibility of the above reaction ? 
 
 (b) Expose a silver coin to the action of hydrogen sulphide ( ?). What 
 happens chemically when silver spoons are used to "beat" an egg? 
 What is "oxidized silver"? 
 
 Silver is "tarnished" by the sulphur in perspiration. 
 
SILVER. 3OI 
 
 Experiment VIII. Action of Sodium Thiosulphate upon Silver Salts. 
 
 Treat small quantities of each of the following substances with a strong 
 cold solution of sodium thiosulphate; silver chloride, silver bromide and 
 silver sulphide (?). Equations? Other insoluble salts of silver behave 
 in a similar manner. 
 
 Sodium thiosulphate ("hyposulphite of soda" or "hypo") is used 
 largely in photograph^: processes for fixing the image; i.e., for dissolving 
 out the unreduced silver halide on the photographic plate. 
 
 Experiment IX. Photography. 
 
 Dip a piece of unsized paper (a filter paper will do) in a sodium chloride 
 solution ; allow the paper to drip for a few seconds, then dip it in a dilute 
 solution of silver nitrate. Place an opaque object upon the paper and 
 expose to the light until the paper around the object has a deep purple 
 color, then dip it successively into a solution of potassium-ferrous oxalate* 
 (a mixture of solutions of ferrous sulphate and potassium oxalate), dis- 
 tilled water, a solution of sodium thiosulphate and distilled water. Dry 
 the paper and fasten it in your laboratory note-book. The foregoing 
 procedure is similar in principle to that involved in the preparation of a 
 "negative" of a picture. Explain the chemistry of each step in the pro- 
 cedure. Equations ? 
 
 Photography involves two processes the preparation of the " negative" 
 and the printing of the "positive." The negative of a picture is usually 
 taken upon a prepared plate of glass, and from this the positive is printed 
 upon sensitized paper. The negative shows the "lights and shades" 
 reversed, hence the name. The positive gives the objects their true 
 appearance, i.e., it is the positive of the original object. 
 
 The modern photographic plate usually consists of a glass plate covered 
 with a gelatin emulsion in which is suspended finely divided silver 
 bromide. This plate is exposed in a camera to the light from the object 
 to be photographed. The silver salt is in part decomposed, the degree 
 of decomposition being proportional to the intensity of the light and the 
 duration of the exposure. Thus a dark object will reflect few light rays, 
 and that part of the plate exposed to these rays will be but little affected. 
 A white surface will act in a reverse manner. 
 
 According to one view, the exposure reduces certain portions of the 
 silver bromide to a sub-bromide, perhaps Ag 2 Br, as indicated by the 
 following equations: 
 
 2 Ag Br Ag 2 Br + Br. 
 
 After exposure, there is usually little or no visible alteration in the film 
 of the plate as the decomposition is only partial. The image must be 
 developed. This is done by washing the plate with a reducing agent, 
 
 *A sodium carbonate solution of pyrogallol is frequently used as the "developer" 
 instead of the potassium-ferrous oxalate. 
 
302 EXPERIMENTAL CHEMISTRY. 
 
 such as an alkaline solution of pyrogallol (photographer's "pyro") or a 
 solution of potassium-ferrous oxalate. This solution, which is known as 
 the developer, continues the action initiated by the light. The partly 
 decomposed silver salt is affected first, and with a speed proportional to the 
 intensity of the illumination undergone by each part. Some of the salt 
 is reduced to metallic silver: 
 
 3 Ag 2 Br + 3FeC 2 4 Fe 2 (C 2 O 4 ) 3 + Fer 3 + Ag, 
 
 The depth of the silver deposit is thus proportional to the intensity of 
 the light upon the original plate. Those portions which have received 
 the most light have the larger amount of the silver salts decomposed and 
 are dark in color. 
 
 It should be remembered that although the developer acts first upon 
 those parts reached by light, it can reduce the whole of the halide upon 
 the plate. When the relative intensities of the light effects i.e., the " con- 
 trast" between the parts variously illuminated have been brought out 
 sufficiently, the plate is removed from the reducing solution. If allowed 
 to remain in the developer long enough, all the silver halide would be re- 
 duced, and the plate would be uniformly dark. 
 
 The plate still contains that portion of the original silver salt which has 
 not been acted upon by the light or the developer. This must be re- 
 moved before the plate is exposed to the light; otherwise, it, too, will be 
 acted upon as previously described, and the image will be blurred or 
 obliterated by more deposits of silver. This is done by a process known 
 as fixing. 
 
 When it is seen by examination that the development has proceeded 
 long enough, the plate is rinsed in water and placed in the fixing bath. 
 This is a solution containing sodium thiosulphate (" sodium hyposulphite " 
 or "hypo"), which is an excellent solvent for many silver compounds. 
 The fixing bath soon dissolves from the gelatin film the silver bromide 
 which remains unaffected by the light or the developer. The " chemistry " 
 of the "fixing" process may be represented by the equation: 
 
 2AgBr + 3Na 2 S 2 O 3 * 2NaAgS 2 O 3 , Na 2 S 2 O 3 * + 2NaBr. 
 
 This treatment is said to "clear" or "fix" the image. The plate is 
 no longer sensitive to light. After thorough washing it is allowed to dry. 
 The result is the negative. 
 
 The plate is now ready to be used in making prints. This is done by 
 allowing sunlight to fall through the negative upon a sheet of sensitized 
 paper, which may be regarded, for purposes of discussion, as paper 
 covered with a film of albumen holding a deposit of silver chloride. The 
 negative and the paper are placed film to film. The action of sunlight 
 upon this paper is similar to that upon the plate. Exposure liberates 
 silver. 
 
 2AgCl->A g2 Cl+_Cl 
 AgCl->Ag+Cl. 
 
 * Formula of the crystals of the complex salt obtained from the solution. 
 
SILVER. 303 
 
 As the dark portions of the negative which correspond to the light 
 parts of the object transmit the fewer light rays (sunlight cannot penetrate 
 the parts over which there is a deposit of silver), it is obvious that what is 
 dark in the negative will be light in the positive and vice versa. The 
 print will have the same shading as the object. 
 
 The print is then toned and fixed. Toning consists in imparting to the 
 print a rich color by replacing part of the Ag 3 Cl with gold from a solu- 
 tion of sodium chlor-aurate, NaAuCl 4 . 
 
 3 Ag 2 Cl + NaAuCl 4 NaCl + 6AgCl + Au. 
 or 
 
 2Ag 2 Cl + K 2 PtCl 4 2KC1 + 4AgCl + Pt. 
 
 Fixing removes the unchanged silver salt, the operation being per- 
 formed as previously described. After thorough washing the picture is 
 dried and mounted. 
 
 "Retouching" is a process whereby blemishes are removed from both 
 the negative and the positive by the use of India ink or colored pencils. 
 
 Experiment X. Silver Nitrate and Organic Matter. 
 
 (a) Press a small crystal of silver nitrate between the thumb and fore- 
 finger for icor 15 seconds. After several hours, examine your finger (?). 
 
 (b) Using a solution of silver nitrate, write your name upon a sheet of 
 paper with a pointed instrument. Expose the writing to the direct sun- 
 light (?). 
 
 What is lunar caustic ? What is the origin of this name ? 
 
 Experiment XI. Precipitation of Silver Chr ornate. 
 
 Add a solution of potassium dichromate to a silver nitrate solution. 
 Describe the result. Equation? 
 
 Experiment XII. Interaction Silver Salts and Potassium Cyanide. 
 
 (a) To a solution of silver nitrate add slowly, drop by drop, a potassium 
 cyanide solution (poison) (?). Equation? Add the precipitant in ex- 
 cess (?). Pass hydrogen sulphide through this solution. Result? 
 Is the silver ion present ? Explain. Equation ?. 
 
 What conclusion with regard to the formation of complex ions may be 
 generally drawn when the precipitate dissolves in excess of the precipi- 
 tant? Give examples. 
 
 (b) All silver salts are soluble in potassium cyanide. Test (?). 
 
 The bath ordinarily used for silver electroplating consists of a solution 
 of potassium argenticyanide. Solutions of the simple salts do not give 
 a coherent film. 
 
 AgNO 3 + 2 KCN <=* KAg(CN) 2 + KNO 3 
 KAg(CN) 2 <=> K- + Ag(CN)' 2 
 Ag(CN)' 2 <=>Ag- 
 Ag' + e <= Ag. 
 
304 EXPERIMENTAL CHEMISTRY. 
 
 Experiment XIII. Analytical Reactions of Silver Salts. 
 
 (a) Compounds of silver heated on charcoal before the blow-pipe give 
 a white metallic globule (?). Test (?). 
 
 (b) Dissolve a small crystal of chemically pure silver nitrate in dis- 
 tilled water. Test the reaction of the solution with litmus paper ( ?). Is 
 the salt hydrolyzed by water ? In this respect it approaches what group 
 of metals? 
 
 What are your conclusions as to the relative chemical activity of silver 
 and copper as metallic elements? 
 
 (c) Those reactions which are of particular importance in analytical 
 chemistry are given in Exps. V and XII. 
 
 (d) Silver chloride is insoluble in either hot or cold water. 
 
 GOLD, All. 
 
 At. Wt. 197.2 Sp. Gr. 19.3. 
 
 Experiment I. Properties of Gold. 
 
 (a) Physical properties. Examine a piece of gold and state briefly 
 the most obvious physical properties of the metal. (The student is re- 
 minded that the "commercial" metal is an alloy of gold and copper. 
 The latter gives it greater hardness). 
 
 Record its melting point, specific heat and atomic heat. Calculate 
 the atomic weight of gold by Dulong and Petit's Law. Pure gold is 
 "24-carat." Explain the use of the word "carat." What is the 
 composition of the American standard gold coin? 
 
 (b) Chemical properties. Is gold affected by free oxygen ? Give 
 reasons for your answer. 
 
 What is the valency of gold? 
 
 Direct a stream of hydrogen sulphide against a piece of pure gold ( ?). 
 The negative result is indicative of its inacitivity. Gold is the least active 
 of all of the familiar metals. 
 
 Does gold displace hydrogen from dilute acids? Before attempting to 
 answer, see table of "Solution Tensions." 
 
 Gold combines with free chlorine, therefore, its solubility in aqua 
 regia with the formation of auric chloride, AuCl 3 . 
 
 Gold does not interact with any of the oxacids except selenic acid. 
 
 Experiment II. Analytical Reactions of Salts of Gold. 
 
 (a) All compounds of gold are decomposed by heat with liberation of 
 the metal. Test, using a solution of auric chloride (?). 
 
 (b) Hydrogen sulphide with solutions of AuCl 3 gives a black precipi- 
 tate, auro-auric sulphide (Au 2 S, Au 2 S 3 ). This sulphide is insoluble in 
 either hydrochloric or nitric acid, but soluble in aqua regia. It is also 
 soluble in ammonium sulphide, forming an ammonium sulpho- or thio- 
 aurate. 
 
 (c) Ferrous sulphate gives a brown or purple precipitate of metallic 
 
GOLD. 305 
 
 gold. If the solution of the auric salt is dilute a bluish color is imparted 
 to the mixture. 
 
 AuCl 3 + 3FeSO 4 ->Fe 2 (SOJ 3 + FeCl 3 + Au. 
 
 (d) Stannous chloride gives with solutions of auric chloride a precipi- 
 tate or coloration which varies in color from reddish-brown to purple. 
 The color is due to the presence of a compound of uncertain composition, 
 known as the purple oj Cassius. 
 
 The test may be made as follows: Add a ferric chloride solution to a 
 solution of stannous chloride until the latter has a permanent yellow color. 
 Add a drop of this solution to an auric chloride solution ( ?), or dip a glass 
 rod into the former solution, and then into the latter (?). 
 
 (e) Other reducing agents, like oxalic acid and potassium nitrite, when 
 warmed with solutions of AuCl 3 , precipitate the gold. If the former re- 
 agent is used, the gold is deposited in a spongy form similar to that used 
 by dentists. 
 
 2 AuCl 3 + 3H 2 C 2 O 4 -> 6HC1 + 6CO 2 + 2 Au. 
 
 2AuCl 3 + 3KNO 2 + 3 H 2 O 3KNO 3 + 6HC1 + 2Au. 
 
 PROBLEMS. 
 
 1. How much cupric oxide is formed by heating 10 grams of copper in 
 the air? 
 
 2. How much cuprous oxide can be prepared, theoretically, by boiling 
 10 grams of cupric sulphate with a solution of potassium hydroxide and 
 glucose ? 
 
 2CuSO 4 .sH 2 O + 4KOH-* Cu 2 O + O + 2K 2 SO 4 + i 2 H 2 O. 
 
 3. Calculate the percentage of copper in malachite, Cu(OH) 2 ,Cu(CO 3 ) 
 In azurite, Cu(OH) 2 ,2CuCO 3 . 
 
 4. If 2 grams of silver chloride yield 1.505 grams of silver, what is the 
 atomic weight of silver? 
 
 5. Calculate the quantities of silver and gold which can be deposited 
 from their respective solutions by the gram-atomic weight of copper. By 
 i gram equivalent. By i gram of copper. 
 
 6. The same electric current passes through separate solutions of silver 
 nitrate and copper sulphate. The electrodes in each case are plates of 
 platinum. The cathode in the copper sulphate solution increased 0.636 
 gram in weight. How much did the cathode in the silver nitrate solution 
 increase in weight? 
 
 7. One coulomb of electricity deposits 0.0011175 gram of silver. This 
 quantity is known as the electro-chemical equivalent of silver. How many 
 coulombs are required for the deposition of one gram equivalent of silver ? 
 Of any substance ? Calculate the electro-chemical equivalent of copper. 
 
CHAPTER XXVI. 
 ALKALINE EARTH METALS. 
 
 (Glucinum, Gl. 9.1) 
 Magnesium, Mg. 24.36 
 Calcium, Ca. 40 . i 
 
 Strontium, Sr. 87 . 6 
 Barium, Ba. 137 .4 
 
 (Radium, Rd. 225.) 
 
 The metals of this family are called the " alkaline-earth" metals because 
 they form the transition from the "alkalies" to the "earth" metals, such 
 as aluminum, gallium, etc. 
 
 None of the elements of the family occur in nature in the uncombined 
 condition, and have only in more recent years been prepared from their 
 compounds. With the exception of magnesium, they are seldom isolated. 
 In the case of glucinum this is due largely to the comparative rarity of 
 its compounds. Calcium, strantium and barium whose compounds 
 are abundant, are seldom prepared in the pure state save in very small 
 quantities because of the extreme difficulty in isolating them, and owing 
 also to the fact that they are of little commercial importance. 
 
 Although these elements are brought into the same family by the peri- 
 odic grouping and possess many marked similarities, yet they do not form 
 an altogether homogeneous and coherent group. 
 
 Glucinum and magnesium are white metals which do not rust (oxidize) 
 rapidly in the air, while calcium, strontium and barium, under similar 
 conditions, quickly lose their silver-white luster. Glucinum does not 
 decompose water; magnesium displaces hydrogen from boiling w r ater; 
 and the other metals decompose water spontaneously at ordinary temper- 
 atures, forming hydroxides. 
 
 Glucinum and magnesium resemble zinc and cadmium and differ from 
 calcium, strontium and barium in that their sulphates are readily soluble 
 in water, but these relations are reversed in the case of their sulphides 
 which are hydrolyzed by water. Glucinum differs from magnesium and 
 resembles zinc in that its hydroxide is acidic as well as basic, i.e., is soluble 
 in sodium or potassium hydroxide. 
 
 Calcium, strontium and barium bear a much closer resemblance to 
 each other in most of their physical and chemical properties than to either 
 glucinum or magnesium. Their hydroxides show an increasing solubility* 
 and alkalinity in the order named. These solutions, which are very 
 dilute, owe their strong alkalinity to the high degree of ionization (see 
 
 * 200, 630, and 2,200 parts of the hydroxides, respectively, dissolve in 1,000,000 
 parts of water. Smith. 
 
 306 
 
ALKALINE EARTH METALS. 307 
 
 table) which they undergo. The order of the solubility of the sulphates 
 of these three metals is reversed 2100, no and 2.3 parts, respectively, 
 being soluble in 1,000,000 parts of water. 
 
 All of these metals form an oxide of the type MO; calcium, strontium 
 and barium also form peroxides of the type MO 2 . The oxides combine 
 with water, forming hydroxides. 
 
 When the nitrates of calcium, strontium, and barium are heated they 
 are decomposed, yielding the oxides of the metals, nitrogen peroxide and 
 oxygen. 
 
 The carbonates which are insoluble, when heated, are decomposed into 
 the oxide of the metal and carbon dioxide. 
 
 The metals of the family usually exhibit a valence of two. 
 
 Radium, one of the recently discovered elements, has never been iso- 
 lated. It is a "radio-active" substance which appears to belong to this 
 family. It exhibits a valence of two in those of its salts which have been 
 investigated. 
 
 GLUCINUM, Gl. 
 
 At. Wt. 9.1 Sp. Gr. 1.8. 
 See lecture-notes or text-book. 
 
 MAGNESIUM, Mg. 
 
 At. Wt. 24.36 Sp. Gr. 1.74. 
 Experiment I. Properties of Magnesium. 
 
 (a) Examine the elementary substance in the forms of "ribbon, " wire 
 and powder. Scrape a piece of magnesium ribbon, and note its color and 
 luster. 
 
 (b) Combustion of magnesium in the air. Introduce a piece of the 
 ribbon into the flame with the forceps ( ?). What is the product ? Equa- 
 tion? 
 
 Apply the flame (caution) to a small pinch of magnesium powder upon 
 an iron plate ( ?). Account for the difference in the speeds of the reactions. 
 
 Note. Powdered magnesium is frequently used as one of the com- 
 ponents of "flash-light powder" for use in photography. 
 
 (c) Solubility in acids. Treat individually small pieces of magnesium 
 ribbon with dilute sulphuric acid ( ?), hydrochloric acid ( ?)> nitric acid ( ?) 
 and acetic acid ( ?). Express each reaction by an equation. 
 
 Experiment II. Magnesium Salts. 
 
 Examine the different salts of magnesium (end shelf), noting the ob- 
 vious physical properties. Give the name, formula, color and solubility* 
 of each salt. Tabulate the foregoing data. 
 
 *See Comey's Dictionary of Chemical Solubilities, and SeidelPs Solubilities of 
 Inorganic and Organic Substances. 
 
308 EXPERIMENTAL CHEMISTRY. 
 
 Experiment III. Hydrolysis of Magnesium Chloride. 
 
 Ascertain by experiment whether magnesium chloride dissolves com- 
 pletely in water (?). Test the solution with litmus paper (?). Explain. 
 Equation ? 
 
 Place several small crystals of the salt in a dry test tube, and heat 
 strongly. Test the reaction of the evolved vapors, and the liquid which 
 condenses in the tube toward litmus paper (?). Explain. Equation? 
 
 Give the formula for carnallite. 
 
 Experiment IV. Precipitation of Magnesium Hydroxide and its Solu- 
 bility in Ammonium Compounds. 
 
 (a) To 10 cm. 3 of solution of magnesium sulphate add a solution of 
 sodium hydroxide or potassium hydroxide (?). Equation? 
 
 (b) Repeat (a) using ammonium hydroxide as the precipitant. Ob- 
 serve the character of the precipitate. Express the reaction by an equa- 
 tion. Is the precipitation of the magnesium complete ? Before attempt- 
 ing to answer this last question, add a strong solution of ammonium 
 chloride (note that this is one of the "products" in the above reaction) 
 in large excess. Explain the disappearance of the precipitate in terms 
 of the ion- or solubility-product constant. 
 
 (c) Can magnesium hydroxide be precipitated in the presence of an ex- 
 cess of an ammonium salt? 
 
 Experiment V. Precipitation of Hydrated Basic Carbonate of Mag- 
 nesium; Its Solubility in Ammonium Compounds. 
 
 (a) Add a slight excess of a sodium carbonate solution to 10 cm. 3 of 
 a solution of magnesium sulphate. The composition of the precipitate 
 varies with conditions. 
 
 (b) Repeat (a), using ammonium carbonate instead of sodium carbonate. 
 Add an excess of a strong solution of ammonium chloride. Result? 
 Equation ? 
 
 (c) Can the basic carbonate of magnesium be precipitated in the pres- 
 ence of an excess of an ammonium salt ? 
 
 Experiment VI. Preparation of Magnesium Chloride. 
 
 To 15 cm. 3 of a dilute solution of hydrochloric acid add an excess of 
 basic magnesium carbonate, filter, evaporate to small volume and crystal- 
 lize. Equation ? What is the formula of the crystallized salt ? 
 
 Experiment VII. Analytical Reactions of Magnesium Salts. Precipi- 
 tation of Ammonium Magnesium Phosphate. 
 
 To 10 cm. 3 of a solution of a soluble magnesium salt add an excess of 
 ammonium chloride, then add cautiously 2 cm. 3 or 3 cm. 3 of ammonium 
 hydroxide or an equal volume of a solution of ammonium carbonate. 
 (Why do neither of the latter substances produce a precipitate?) To 
 the clear solution thus obtained add a disodium hydrogen phosphate 
 
ALKALINE EARTH METALS. 309 
 
 (Na 2 HPO 4 ) solution until precipitation ceases. Describe the precipitate. 
 Express the reaction by an equation. 
 
 When ammonium magnesium phosphate is ignited it is converted into 
 the anhydrous pyrophosphate of magnesium. 
 
 ( 2 NH 4 MgP0 4 , 6H 2 -> Mg 2 P 2 7 + 2 NH 3 + 13^0). 
 What is the color and valence of the magnesium ion ? 
 
 CALCIUM, ca. 
 
 At. Wt. 40.1 Sp. Gr. 1.85. 
 Experiment I. Properties of Calcium. 
 
 (a) Enumerate the most obvious physical properties of the metal. 
 
 (b) Does it interact with water? Equation? 
 
 (c) Will it burn in the air? Equation? 
 
 Note. The product of the combustion is a mixture of the oxide and 
 the nitrite (Ca 3 N 2 ). Compare with the combustion of magnesium in air. 
 
 Experiment II. Action of Acids upon Calcium Carbonate. 
 
 (a) Recall or repeat Exp. II (c), "Carbon." Indicate by an equation 
 the action of hydrochloric acid upon marble (impure calcium carbonate). 
 
 (b) Test a small quantity of powdered marble with dilute sulphuric 
 acid, and note the effect. Equation ? 
 
 (c) Repeat (b) substituting nitric acid for sulphuric acid. 
 
 Experiment III. Preparation of Calcium Oxide (Quicklime). 
 
 Ignite about 3 grams of powdered marble for 15 to 20 minutes in an 
 open crucible placed upon a pipe-stem triangle. Maintain the crucible 
 and contents at the highest temperature obtainable with the aid of a blast- 
 lamp. Occasionally stir the mass with a platinum wire. When the mass 
 has cooled, add about 10 cm. 3 of water; stir the mixture (?). Test reac- 
 tion of the liquid with litmus paper (?). Does water act upon marble? 
 Add a few drops of hydrochloric acid to the above mixture (?). How 
 do the results compare with the action of acid upon marble ? What are 
 your conclusions as to the identity of the substance in the crucible ? The 
 decomposition of the marble by heat is a reversible reaction. How may 
 it be brought to a condition of equilibrium? How may it be reversed? 
 How may it be reversed without altering the temperature? What pre- 
 cautions must be taken in the construction of a lime-kiln that the complete 
 decomposition of marble into calcium oxide and carbon dioxide may be 
 insured? Indicate by equations the chemical equilibria involved in the 
 above decomposition. 
 
 Experiment IV. Preparation and Properties of Calcium Hydroxide. 
 Lime-Water. 
 
 (a) Pour 5 cm. 3 of distilled water upon 8 grams of quicklime in a porce- 
 
310 EXPERIMENTAL CHEMISTRY. 
 
 lain dish, and note the results (?). The foregoing process is known as 
 the "slaking" or "slacking" of lime. The product is known as calcium 
 hydroxide. Express the reaction by an equation. What is "mortar"? 
 What is "air-slacked" lime? 
 
 (b) Solubility* and effect of calcium hydroxide upon litmus. Put 
 4 or 5 grams of the calcium hydroxide prepared in (a) into a flask or 
 bottle containing 200 cm. 3 of distilled water, cork tightly, shake vigorously 
 from time to time, and allow the mixture to stand until the solution has 
 become clear, when the latter may be decanted, drawn off with a siphon or 
 removed with a pipette. If too much time is consumed in allowing the 
 mixture to settle, it may be filtered rapidly, and the clear liquid (lime- 
 water) used for the following experiments. 
 
 (c) Test the action of the solution upon litmus paper (?). The result 
 indicates the presence of what ion ? Explain. 
 
 (d) Recall or repeat Exp. II (d), "Carbon." Equations? 
 
 To a piece of old mortar in a test tube add dilute hydrochloric acid. 
 Identify the gas (?). What chemical action is involved in the "harden- 
 ing" of mortar? 
 
 (e) Recall or repeat Exp. V "Carbon." Equations? When is water 
 said to possess " temporary hardness" ? It is due to the presence of what 
 salt? 
 
 Experiment V. Calcium Salts. 
 
 Repeat Exp. II "Magnesium," substituting the word calcium where 
 the word magnesium appears. 
 
 Experiment VI. Preparation of Calcium Salts. 
 
 (a) Calcium chloride. Suggest by an equation such a method for 
 the preparation of calcium chloride from one of its salts as will not in- 
 volve the simultaneous formation of any other salt, acid or base. Is 
 calcium chloride deliquescent or efflorescent? It is frequently used for 
 what purpose in connection with the manipulation of gases? 
 
 (b) Calcium hypochlorite. Describe a method for the preparation of 
 this salt. (Suggestion, see Exp. Ill (a), " Chlorine.") Equation ? What 
 is "bleaching powder "?f Its formula? 
 
 Experiment VII. Dehydration of Gypsum. Plaster of Paris. 
 
 (a) Heat a few grams of powdered gypsum or a small piece of a selenite 
 crystal (CaSO 4 . 2H 2 O) in a test tube and record the obvious results. 
 Pulverize the residue and ascertain whether it will become solid when 
 
 * At 1 8, 600 parts of water by weight dissolve i part of the hydroxide; at 100, 
 about twice as much water is required to dissolve the same quantity of salt. Seidtll. 
 
 fThe probable formula of bleaching powder is CaCl (CIO) a mixed salt rather 
 than an equi-molar mixture of CaCl 2 andCa(OCl) 2 . This view is supported by the 
 fact that CaCl 2 cannot be dissolved out with alcohol, nor is the salt deliquescent as is 
 calcium chloride (CaCl 2 ). The ions Ca", Cl', and CIO' are present in a solution of 
 the salt. 
 
ALKALINE EARTH METALS. 311 
 
 mixed with a little water to form a paste and allowed to stand (?). See 
 whether gypsum itself will act in the same way with water ( ?) . Equations ? 
 
 Note. If all of the water is removed from gypsum by heating or the 
 temperature is allowed to rise much above 125, the product when mixed 
 with water does not "set" quickly. 
 
 (b) Make a thick paste by mixing a little water with "plaster of Paris." 
 Place the paste upon a glass plate and make a cast of a coin. Chemically, 
 what is "plaster of Paris?" Probable formula? 
 
 Experiment VIII. Solubility of Phosphates of Calcium. 
 
 (a) Ascertain whether or not the tertiary orthophosphate of calcium 
 (Ca 3 (PO 4 ) 2 ) is soluble in water? Would such a salt make a good 
 fertilizer ? Reasons for your answer ? 
 
 (b) Repeat (a) using primary calcium phosphate (Ca(H 2 PO 4 ) 2 . 
 What is "superphosphate of lime"? Its use? 
 
 Experiment IX. Analytical Reactions of Calcium Salts. 
 
 (a) Flame test. Introduce a small quantity of calcium chloride into 
 the Bunsen flame by means of a platinum wire (?). Examine the flame 
 with a spectroscope, noting particularly the presence of two bands a 
 red and a green one which impart the brick-red color to the flame. 
 
 (b) (Use a dilute solution of pure calcium chloride for the following 
 tests.) Examine such solutions of calcium salts as are found on the end 
 shelf. What is the color of the calcion? Its valence? Justify your 
 conclusions. 
 
 (c) Calcium carbonate. To a portion of the solution containing the 
 calcion add a slight excess of an ammonium carbonate solution (?). 
 Warm the solution if necessary to procure complete precipitation. Ex- 
 press the reaction by an equation. Filter, spread the filter paper with 
 precipitate upon a glass plate and divide the precipitate into two parts. 
 Treat one part with hydrochloric acid ( ?) and the other with dilute 
 acetic acid (?). Equations? What are your conclusions as to the rela- 
 tive strength of acetic and carbonic acids? Confirm your inference by 
 giving the degrees of ionization in .1 normal solutions of the acids. Ex- 
 plain in terms of the "electrolytic dissociation theory" how the insoluble 
 salt of a "weak" acid is dissolved by a stronger acid. 
 
 (d) Calcium oxalate. Add an excess of an ammonium oxalate solu- 
 tion to a dilute solution of calcium chloride (?). Equation? Filter and 
 divide precipitate into two parts. Test the precipitate as in (c). Inter- 
 pret results. 
 
 Explain the difference in behavoir of the oxalate and carbonate of 
 calcium toward acetic acid, taking into account the difference in solubility 
 of the two salts, and the behavior of oxalic acid and carbonic acid, re- 
 spectively. Why was ammonium oxalate used as the preciptant in pref- 
 erence to oxalic acid (recall or repeat Exp. II (c~), "Chemical Equilib- 
 
312 EXPERIMENTAL CHEMISTRY. 
 
 rium")? Under what conditions may calcium oxalate be precipitated 
 from a calcium chloride solution ? Explain. 
 
 (e) I. Calcium sulphate. (e,I. Repeat Exp. II (ft), "Chemical Equi- 
 librium.") 
 
 (e) II. To a dilute solution of calcium chloride add a slight excess 
 of dilute sulphuric acid ( ?) Equation ? Filter and divide the nitrate 
 into two parts. Add sufficient ammonium hydroxide to one portion to 
 barely neutralize the free acids (determine the "end point" by means 
 of litmus paper). Now add ammonium oxalate to the neutralized solu- 
 tion (?). Explain the precipitation of calcium oxalate. Which is the 
 more soluble the sulphate or oxalate of calcium ? Give the actual solu- 
 bilities of each. (See works to which you have been referred previously.) 
 
 (e) III. "Permanent hardness" of water. Boil the second portion of 
 the nitrate from (e), II. Compare the negative result with that obtained 
 in Exp. IV (e). To the above solution add a solution of sodium carbo- 
 nate (?). Equation? Devise a method for proving that the precipitate 
 is calcium carbonate (?). Is the sulphate or the carbonate more solu- 
 ble (?). Reasons for your answer? Give the actual solubilities of the 
 two substances. Can "permanent hardness" be removed by boiling? 
 Give one method for removing it. What are some of the objections to 
 "hard water" for domestic and technical use? 
 
 STRONTIUM, Sr. 
 
 At. Wt. 87.6 Sp. Gr. 2.54. 
 Experiment I. Properties of Strontium and its Salts. 
 
 Strontium is a yellowish-white, rather tough metal which slowly oxi- 
 dizes spontaneously in the air. It burns when heated in the air, and de- 
 composes water at ordinary temperatures. It is malleable and ductile and 
 fuses at red heat. 
 
 The physical and chemical properties of the compounds of strontium are 
 very similar to those of calcium. 
 
 Experiment II. Strontium Salts. 
 
 Repeat Exp. II, "Magnesium," substituting the word strontium for 
 magnesium. 
 
 Experiment III. Analytical Reactions of Strontium Salts. 
 
 (a) Flame test. Dip a platinum loop into a solution of strontium 
 chloride and hold it in the Bunsen flame ( ?). Examine the flame with a 
 spectroscope and make a sketch of the spectrum showing the position of 
 the lines with reference to the sodium lines. 
 
 Note. Anhydrous strontium is mixed with charcoal, sulphur and potas- 
 sium chlorate to make "red fire." 
 
 (b) Strontium carbonate. Add ammonium hydroxide and ammonium 
 
ALKALINE EARTH METALS. 313 
 
 carbonate to a dilute solution of strontium chloride (?). Equation? 
 Filter, wash the precipitate and test the action of acetic and hydrochloric 
 acids upon it (?). Equations? 
 
 (c) Strontium sulphate. To a solution containing strontium add a 
 small amount of a solution of calcium sulphate; heat to boiling, and if no 
 precipitate appears immediately, let the mixture stand for 10 or 15 minutes. 
 Explain why the precipitate forms so slowly. What do you infer regard- 
 ing the relative solubilities of calcium and strontium sulphates. Give 
 the actual solubilities of each. 
 
 What is the color of the strontium ion ? Its valence ? 
 
 BARIUM, Ba. 
 
 At. Wt. 137.4 Sp. Gr. 3.6. 
 Experiment I. Properties of Barium and its Salts. 
 
 Barium is a bright yellow metal which readily oxidizes spontaneously 
 in the air. It decomposes water energetically at ordinary temperatures. 
 The metal is malleable and ductile and fuses at red heat. 
 
 The physical and chemical properties of the compounds of barium 
 resemble those of calcium and strontium. 
 
 Experiment II. Barium Salts. 
 
 Same as Exp. II, "Strontium. " 
 
 Experiment III. Water of Hydration in Barium Chloride. 
 
 A crucible with the lid, which has been cleaned, is placed on a pipe- 
 stem triangle and heated with the Bunsen burner for a few minutes. 
 Cool, and determine carefully the weight of the crucible. Weigh into 
 crucible accurately about 3 grams of pure crystallized barium chloride. 
 Place the crucible with its contents in an air bath at 120 to 130 C. for 
 about one hour, then cool for 10 minutes and weigh; or the crucible and its 
 contents may be placed on a triangle and heated gently with a small Bunsen 
 burner flame. The temperature is gradually raised until the crucible 
 attains a low red heat at which it is maintained for 10 minutes when it is 
 cooled and weighed as per above directions. The heating and weighing 
 is repeated until the weight is constant. 
 
 Record your results as follows: Grams. 
 
 Weight of Xble + barium chloride = 
 Weight of Xble alone = 
 
 Weight of barium chloride taken = 
 Weight of Xble + contents before heating = 
 Weight of Xble + contents after heating = 
 
 Weight of water found = 
 Per cent, of water in barium chloride = 
 Theoretical per cent. = 
 
 Error = 
 
314 EXPERIMENTAL CHEMISTRY. 
 
 Experiment IV. Analytical Reactions of Barium Salts. 
 
 (a) Flame test. Same as Exp. Ill (a), "Strontium." 
 
 Note: Barium chlorate is used with charcoal and sulphur to prepare 
 "green fire." 
 
 (b) Barium carbonate. Same as Exp. Ill, b, "Strontium." 
 
 (c) Barium chromate. To a dilute solution of barium chloride add a 
 solution of potassium chromate (?). Ascertain whether the precipitate 
 is soluble in dilute acetic acid. In hydrochloric acid. 
 
 Note. Neither calcium nor strontium chromate is precipitated from 
 dilute solutions acidified with acetic acid. 
 
 (d) Barium sulphate. Add dilute sulphuric acid to a barium chloride 
 solution (?). Equation? 
 
 To 10 cm. 3 of a solution of barium chloride add 20 cm. 3 of a clear 
 saturated solution of strontium sulphate (prepared by shaking the salt 
 with distilled water and filtering) (?). Which is the more soluble stron- 
 tium or barium sulphate? 
 
 Arrange the sulphates of calcium, strontium and barium in the order 
 of decreasing solubility. Give the actual solubility of each. 
 
 Experiment V. Detection of the Alkaline Earth Metals in a Mixture. 
 
 (a) Give two methods for distinguishing the compounds of the metals 
 of this group. 
 
 (b) How may the carbonate of magnesium be separated from the car- 
 bonates of the other three metals of this group ? 
 
 Are the carbonates of calcium, strontium and barium soluble in acetic 
 acid? 
 
 State how barium may be separated from calcium and strontium. 
 How may the presence of calcium and strontium in a solution of their 
 compounds be proven? 
 
 (c) If a solution of the compounds of the four metals of this group were 
 given you, how would you proceed to prove their presence? Make a 
 statement of your proposed method and submit it to the instructor for 
 inspection. 
 
 (d) Procure an "unknown" solution containing two or more of the 
 elements of this group and make an analysis of same according to the 
 method proposed in (c). 
 
 RADIUM, Rd. 
 
 At. Wt. 225 Sp. Gr. (?) 
 (See lecture notes and references.) 
 
 EXERCISES. 
 
 i. Chemically, what is (a) asbestos? (b) meerschaum? (c) olivine? 
 (d) serpentine? 
 
ALKALINE EARTH METALS. 315 
 
 2. What is (a) chalk? (b) calcite? (c) limestone? (d) cement 
 (Portland) ? (e) apatite ? 
 
 3. What is (a) celestite? (b) strontianite ? 
 
 4. What is (a) witherite ? (b) heavy spar or barite ? 
 
 PROBLEMS. 
 
 1. Calculate the percentage of barium in barium chloride. 
 
 2. Calculate the weight of sulphuric acid necessary to precipitate the 
 barium as barium sulphate from a solution containing .5 gram of barium 
 chloride, BaCl 2 . 2 H 2 O? 
 
 3. Calculate the percentage of chlorine in barium chloride. 
 
 4. Calculate the weight of silver nitrate required to precipitate the 
 chlorine as silver chloride from a solution containing .7 gram of barium 
 chloride. 
 
CHAPTER XXVII. 
 
 Zinc, Zn, 65.4 
 
 Cadmium, Cd, 112.4 
 Mercury, Hg, 200.0 
 
 Zinc, cadmium and mercury form the secondary family of the alkaline 
 earths. The elementary substances of this family and their corre- 
 sponding compounds possess many similar properties. Their relation 
 to magnesium, calcium, strontium and barium resembles the relation- 
 ship existing between the alkalies and copper, silver, and gold. 
 
 These three elementary substances are lustrous metals of high specific 
 gravity and are found in combination as minerals, which occur usually 
 as ores in the older crystalline rocks. Smithsonite, ZnCO 3 , sphalerite 
 or zinc blende, ZnS, and calamine, H 2 Zn 2 SiO 5 , are the principal minerals 
 which contain zinc. Cadmium is found in small quantities in those 
 ores in which zinc is the chief metal present. It is also found in the 
 rare mineral, greenockite, CdS. Small quantities of mercury are found 
 in minute globules disseminated through the pores of those rocks where 
 the most important mineral containing mercury, namely, the red sulphide, 
 cinnabar (HgS), occurs. Alloys of silver and gold, known as amalgams, 
 are also found native. The metallurgy of these substances is comparatively 
 simple all three of the metals being obtained easily by roasting the ores 
 alone or in mixture with carbon, in ovens, and condensing the vapors 
 as the metals volatilize at high temperatures. 
 
 In their physical characteristics, the elements of this family are cer- 
 tainly metallic in their nature. Zinc, which has the smallest atomic 
 weight of the three, is less positive than the other two. The appended 
 table shows the interesting gradation of properties, physical and chemical, 
 with increasing atomic weight; for example, the boiling points and the 
 melting points decrease with increase of atomic weight. It will be 
 remembered that the non-metallic elements behave in a contrary manner 
 and that the conduct of the alkali metals is very similar. 
 
 Zinc is a bluish-white crystalline metal, brittle at ordinary tempera- 
 tures, but malleable at i2o-i5o. Cadmium is a white, lustrous metal, 
 fairly soft, ductile and malleable. Mercury is a heavy, silvery, lustrous 
 liquid. 
 
 At ordinary temperatures zinc is acted upon but superficially by moist 
 air, but cadmium is scarcely affected and mercury not at all. However, 
 all three of the metals are converted into oxides when heated in the air. 
 
 Finely divided zinc will set free hydrogen from water at ordinary 
 temperature, but massive zinc and cadmium act upon water only at red 
 heat. Mercury has no effect upon water. 
 
ZINC. 317 
 
 The metals of this family can be deposited from solutions of their 
 salts by the electric current or by metals above them in the electro- 
 chemical series, and they will displace other metals .below them in the 
 series. 
 
 These metals are bivalent mercury being the only one which forms 
 two series of salts (mercurous and mercuric compounds), which differ 
 from each other in the proportions in which mercury is contained in 
 them. The same tendency is shown to a slight extent by cadmium. 
 
 Zinc dissolves easily in dilute acids, and is attacked by the caustic 
 alkalies when the latter are warmed. Hydrogen is evolved in both cases. 
 Cadmium dissolves less readily in acids, but is not soluble in solutions 
 of the caustic alkalies. Mercury does not dissolve in hydrochloric acid 
 or dilute sulphuric acid; concentrated sulphuric acid, when hot, however, 
 attacks it and liberates sulphur dioxide; concentrated nitric acid acts 
 upon the metal forming mercuric nitrate; dilute nitric acid yields mer- 
 curous nitrate. The caustic alkalies do not attack mercury. 
 
 The oxides and hydroxides of zinc, cadmium and mercury are soluble 
 in acids, yielding salts, in solutions of which the metals are positive 
 ions. A few zinc salts, known as zincates and of the type Na 2 ZnO 2 , in 
 which zinc is in the negative ion, are formed by the interaction of zinc, 
 its oxide or hydroxide, with strong basic hydroxides. 
 
 2NaOH + Zn > Na-jZnO-s +~H 2 
 2 NaOH + Zn(OH) 2 Na 2 ZnO 2 + 2 H 2 O. 
 
 In this respect zinc resembles aluminum, lead and a number of other 
 metals which can display both metallic and non-metallic properties. 
 
 The chlorides, nitrates and sulphates are soluble in water, but the 
 oxides, carbonates and phosphates are insoluble in water. The chlorides 
 are comparatively volatile and the hydroxides lose water forming the 
 oxides. 
 
 The compounds of zinc, cadmium and mercury do not give color to 
 the borax bead. 
 
 The following table emphasizes the similarity in the properties of the 
 three metals and their corresponding derivatives. 
 
 Zinc. Cadmium. Mercury. 
 
 Atomic Weight, 65.4 112.4 200.0 
 
 Specific Gravity, 7.1 8.6 13-59, (o) 
 
 Atomic Volume, 9.1 12.9 14.7 
 
 Melting-point, 419 321.7 38.85 
 
 Boiling-point, 920 (?) 778 357 
 
 *Vapor Density, 65.5 112 200.3 
 
 * These vapor density determinations indicate that the molecular weight and the 
 atomic weight of each of the. three elementary substances, when they are in the state 
 of vapor, are identical. If our conclusions, as to the relative weights of these atoms, 
 are correct, it is obvious that zinc, cadmium and mercury, on changing from the 
 liquid to the gaseous state, separate into individual atoms. Sodium and potassium 
 behave in a similar manner. 
 
ZnO,H 2 
 82,600 cal. 
 
 CdO,H 2 O 
 65,700 cal. 
 
 HgO 
 
 22,000 cal 
 
 ZnCl 2 
 97,200 cal. 
 
 CdCl 2 
 93,200 cal. 
 
 HgCl 2 
 
 54,500 cal. 
 
 ZnS 
 (Zn,S,Aq), 
 39,600 cal. 
 
 CdS 
 (Cd, S, Aq.), 
 32,400 cal. 
 
 HgS 
 Hg,S, 
 16,200 cal. 
 
 ZnSO 4 
 230.000 cal. 
 
 CdSO 4 
 221.000 cal. 
 
 HgS0 4 . 
 
 318 EXPERIMENTAL CHEMISTRY. 
 
 Oxides, 
 H. of F. 
 
 Chlorides, 
 H. of F. 
 
 Sulphides, 
 H. of F. 
 
 Sulphates, 
 H. of F. 
 
 Note. It is maintained by many that if we were able to make vapor 
 density determinations at very high temperatures, we should discover 
 that all diatomic molecules are converted into monatomic ones. 
 
 ZINC, zn. 
 
 At. Wt. 65.4 Sp. Gr. 7.1. 
 Experiment I. Properties of Zinc. 
 
 (a) Clean part of a piece of zinc with sand-paper or a file. Color ? 
 
 (b) Is zinc hard or soft (use the point of a knife blade) ? 
 
 (c) Does zinc tarnish when exposed to the moist air ? 
 
 (d) Is zinc soluble in sulphuric acid ? In hydrochloric acid ? Equa- 
 tions? 
 
 (e) Place about i cm. 3 of zinc dust in a test tube, and add 5 cm. 3 of a 
 strong solution of sodium hydroxide. Heat the mixture carefully and 
 test the evolved gas with a flame. Results ? Identify the gas. The 
 solution contains sodium zincate, Na 2 ZnO 2 . Express the action by an 
 equation. 
 
 (/) Recall or repeat Exp. XI (a), " Electrolysis and Electrical Equiva- 
 lents." 
 
 (g) Enumerate some of the uses of zinc. What is "galvanized" 
 iron? 
 
 Experiment II. (Quant.) Determination of the Atomic Weight of 
 Zinc. 
 
 (a) Proceed as in Exp. XIII, "Hydrogen." Calculate the weight of 
 hydrogen collected. Determine the atomic weight of zinc referred to 
 H = i, remembering that zinc is bivalent. Write the reaction. 
 
 (b) The specific heat of zinc is 0.093. Calculate the atomic weight 
 of zinc by Dulong and Petit's Law. State the latter. 
 
 Experiment III. Preparation of Hydrated Zinc Chloride. 
 
 Dissolve two or three small pieces of "granulated" zinc in hydro- 
 chloric acid, dilute with an equal volume of water, filter and evaporate 
 
ZINC. 319 
 
 the filtrate to oily consistency in a porcelain dish. Cool, and examine 
 the product (?). Equation? 
 
 Experiment IV. Preparation of Hydrated and Dehydrated Zinc Sul- 
 phate. 
 
 (a) Suggest a method for the preparation of hydrated zinc sulphate, 
 ZnSO 4 . yH 2 O. Equation ? 
 
 (b) How may the above salt be dehydrated? At what temperature 
 does it lose all or a part of its water of hydration. (Consult your text- 
 book) ? What is white vitnol? 
 
 Experiment V. Precipitation and Behavior of Zinc Hydroxide. Form- 
 ation of a Zincate. 
 
 (a) To a dilute solution of zinc sulphate add a sodium hydroxide 
 solution drop by drop and shake. Result? What, probably, is the 
 precipitate? Equation? Filter, and suspend the precipitate in water 
 in a test tube by punching a hole through the apex of the filter paper, 
 and directing a stream of water from the water-bottle upon the pre- 
 cipitate. Divide the mixture into two parts. 
 
 (b) To one part add hydrochloric acid? Equation? Does the zinc 
 hydroxide behave in this reaction as a base or an acid ? 
 
 (c) To the second part of the mixture from (a) add an excess of the 
 sodium hydroxide solution (?). The compound formed by the inter- 
 action of zinc hydroxide and sodium hydroxide is sodium zincate, 
 Na 2 ZnO 2 . Ionic equation? Does this reaction indicate that zinc hy- 
 droxide possesses acidic or basic properties? Is the zinc a part of the 
 anion or cation ? 
 
 What are your inferences as to the "strength" of zinc hydroxide as a 
 base ? 
 
 Experiment VI. Preparation and Properties of Zinc Hydroxycarbon- 
 ate and Zinc Oxide. 
 
 (a) Add slowly a solution of sodium carbonate to a dilute solution of 
 zinc sulphate ( ?). Note the evolution of gas as the precipitate forms. 
 Prove that the escaping gas is carbon dioxide (?). Boil the mixture, 
 taking care that the precipitant is in excess; filter, and wash the pre- 
 cipitate. Test a small portion of the precipitate with hydrochloric acid. 
 Is carbon dioxide evolved ? Account for the evolution of this gas during 
 the formation of the precipitate. Express by ionic equations the pre- 
 cipitation of the basic zinc carbonate. 
 
 (b) Heat the other portion of the precipitate from (a) to redness in a 
 porcelain crucible until a portion removed and tested with hydrochloric 
 acid does not effervesce. Note the color of the residue when hot and 
 when cold (?). Equation? Reserve the zinc oxide for Exp. VIII (a). 
 
320 EXPERIMENTAL CHEMISTRY. 
 
 Experiment VII. An Experimental Study of Ionic Equilibrium and 
 "Concentration Effect." 
 
 (a) Place 5 cm. 3 of zinc acetate, zinc sulphate and zinc chloride in 
 separate test tubes. Test each solution with litmus paper ( ?). Saturate 
 the solutions with hydrogen sulphide. Results? Equations? Are 
 these reactions reversible theoretically; i.e., does the acid which is formed 
 simultaneously with the zinc sulphide in each case dissolve the latter to a 
 larger or smaller degree? Will the amount of zinc sulphide dissolved 
 vary with the activity of the different acids? The answers to these 
 latter questions may be ascertained by two different experimental methods. 
 Proceed as follows: 
 
 Filter separately the contents of the tubes and preserve each filtrate 
 in a separate tube which should be marked to prevent confusion. 
 
 Compare the rates at which zinc sulphide (the precipitate) dissolves 
 in dilute acetic acid, dilute sulphuric acid and dilute hydrochloric acid. 
 Results? What are your inferences as to the relative "activities" of the 
 three acids? 
 
 Confirm your inferences by adding ammonium hydroxide to each 
 filtrate in the marked tubes. The zinc in solution is now precipitated 
 as zinc sulphide (the ammonium hydroxide merely neutralizes the free 
 acid). In which solution (filtrate) do you find the largest amount of 
 zinc ? The least amount ? Do these results confirm your previous 
 conclusions as to the relative strength (activity) of the three acids ? 
 
 (b) To 5 cm. 3 of a dilute solution of zinc sulphate add a few drops 
 of sulphuric acid. Pass hydrogen sulphide into the solution. If a 
 precipitate forms, add more acid, drop by drop, until hydrogen sulphide 
 does not produce a precipitate. The addition of sulphuric acid 
 increases the concentration of what ions? Does this influence the 
 degree of ionization of the hydrogen sulphide ? When there is an excess 
 of free acid present in the solution is the hydrogen sulphide chiefly in 
 the molecular condition or is it largely dissociated ? Is chemical action 
 an interaction of ions or molecules ordinarily ? Show by ionic equation 
 the failure of hydrogen sulphide to precipitate zinc sulphide in the presence 
 of a free inorganic acid. Now add an excess of a sodium acetate solution 
 (zinc acetate is soluble), and pass hydrogen sulphide into the solution 
 if a precipitate of zinc sulphide does not form at once. Account for the 
 formation of the precipitate. Express the various interactions by ionic 
 equations. 
 
 Experiment VIII. Analytical Reactions of Zinc Salts. 
 
 (a) Heat a small piece of zinc on charcoal in the oxidizing flame ( ?). 
 Moisten the incrustation which is formed on the charcoal with a drop of 
 cobalt nitrate, and heat again or place a small pinch of the zinc oxide 
 prepared in Exp. VI (b) on the charcoal and proceed as above. Result ? 
 Equations? (The formula of the green compound, Rinmaris green, is 
 CoZnO 2 .) 
 
CADMIUM. 321 
 
 Does the zinc manifest acidic or basic properties in this reaction? 
 Reasons for your answer? 
 
 (b) Test a dilute solution of zinc sulphate with litmus paper (?). 
 Account for the result. 
 
 (c) Ammonium hydroxide produces in neutral solutions of zinc salts 
 a partial precipitation of zinc hydroxide which is soluble in excess 
 of the precipitant and ammonium chloride, forming Zn(NH 3 ) 4 .Cl2 or 
 Z n(NH 3 ). 4 (OH) 2 . If this solution is treated with ammonium sulphide, 
 (NH 4 ) 2 S, zinc sulphide is formed. Equations? 
 
 (d) Add a slight excess of ammonium sulphide to several cm. 3 of a 
 solution of zinc sulphate. Result ? Equation ? Filter, and wash the 
 precipitate with water. Divide the precipitate into two parts. Try 
 its solubility in dilute hydrochloric acid, and acetic acid ( ?). Equations? 
 
 Why was ammonium sulphide used instead of hydrogen sulphide 
 in the above reaction ? 
 
 What is the color of the zinc ion ? Its valence ? 
 
 CADMIUM, Cd. 
 
 At. Wt. 112.4 Sp. Gr. 8.6. 
 
 Experiment I. Properties of Cadmium. 
 
 (a) Examine a piece of the metal and tabulate its obvious physical 
 and chemical properties. 
 
 (b) Will cadmium displace zinc from a solution of a salt of the latter 
 (Note their relative positions in the electro-chemical series) ? 
 
 (c) Calculate the atomic weight of cadmium by Dulong and Petit's 
 Law, the specific heat of cadmium being 0.054. 
 
 (d) Give the composition of "Wood's Metal." (See Appendix). Re- 
 peat or review Exp. II. "Bismuth." 
 
 (e) Enumerate some of the chief uses of the metal. 
 
 Experiment II. Properties of Cadmium Salts. 
 
 Same as Exp. II "Magnesium." 
 
 Experiment III. Precipitation and Solubility of Cadmium Hydroxide. 
 
 Repeat Exp. V (a) and (b), substituting the word cadmium where 
 the word zinc appears. Equations? Compare the behavior of the zinc 
 and the cadimum compounds (?). 
 
 Experiment IV. Precipitation of Cadmium Carbonate. Preparation 
 of Cadmium Oxide. 
 
 (a) To 10 cm. 3 of a dilute solution of cadmium chloride add a small 
 excess of sodium carbonate. Result ? Equation ? Filter, and wash 
 the precipitate. Test a small portion of the precipitate with hydrochloric 
 acid (?). Equation? 
 
 (b) Heat the rest of the precipitate from (a) to redness in a crucible 
 
322 EXPERIMENTAL CHEMISTRY. 
 
 until a portion removed and tested with hydrochloric acid does not 
 effervesce. Note the color of the residue when hot and when cold (?). 
 Equations? Reserve the cadmium oxide for Exp. V (a). 
 
 Experiment V. Analytical Reactions of Cadmium Salts. 
 
 (a) Mix a portion of the cadmium oxide obtained from Exp. IV (b), or 
 some other cadmium compound, with anhydrous sodium carbonate, 
 and heat the mixture on a piece of charcoal in the reducing flame. Dur- 
 ing the preliminary stages of the reduction, note the characteristic in- 
 crustation (?). Continue the process until small metallic globules (?) 
 make their appearance. 
 
 (b) Same as Exp. VIII (b) "Zinc." 
 
 (c) Ammonium hydroxide precipitates cadmium hydroxide from 
 solutions of cadmium salts. The hydroxide is soluble in excess, form- 
 ing salts such as Cd(NH 3 ) 4 .SO 4 or Cd(NH 3 ) 4 . (OH) 2 . If this ammo- 
 niacal solution is treated with a solution of potassium cyanide, KCN, a 
 soluble complex salt, K 2 .Cd(CN) 4 is formed. 
 
 (d) Saturate a solution of cadmium chloride with hydrogen sulphide ( ?) . 
 Equation? Is the reaction reversible? Answer this question 
 by testing the solubility of cadmium sulphide in dilute hydrochloric acid. 
 Will hydrogen sulphide precipitate cadmium sulphide from solutions 
 of cadmium salts containing a slight amount of free acid? Recall the 
 behavior of zinc salts in the presence of the same reagents ( ?). 
 
 Test the solubility of cadmium sulphide in concentrated nitric acid ( ?). 
 Equation? In ammonium sulphide (?). 
 
 (e) To a solution of cadmium chloride add a solution of potassium 
 cyanide (Care! Poison!), drop by drop. Observe the precipitation of 
 cadmium cyanide, Cd(CN) 2 . Equation? Add an excess of the pre- 
 cipitant (?). The soluble complex salt formed is K 2 .Cd(CN) 4 . Equa- 
 tion? Is the cadmium ion present in the solution, theoretically? (See 
 discussion of "double" and " complex" salts under Exp. XVII, 
 "Copper".) Confirm or disprove your conclusions by passing hydrogen 
 sulphide into the above solution. Results? Ionic equations? Is 
 cadmium sulphide soluble in potassium cyanide ? 
 
 (/) What is the color and valence of the cadmium ion ? 
 
 (g) State the tests by which you could distinguish between the salts 
 of zinc and cadmium. Apply to the assistant for an " unknown " solution. 
 Ascertain if the solution contains one or both of the above metals. Make 
 a complete report of your procedure. 
 
 MERCURY, Hg. 
 
 At. Wt. 200.0 Sp. Gr. 13.59. 
 Experiment I. Properties of Mercury. 
 
 (a) Tabulate the most obvious physical properties of mercury. What 
 particular property does it possess which makes it unique among metals ? 
 
MERCURY. 323 
 
 (b) Does mercury tarnish when exposed to the air? 
 
 (c) Enumerate those properties of mercury in virtue of which it appeals 
 to you as a metal. 
 
 (d) Is the molecule of mercury monatomic or diatomic ? 
 
 (e) State some of the uses of mercury. What are amalgams ? 
 
 Experiment II. Preparation of Mercury by Roasting Cinnabar. 
 
 Place a small piece of cinnabar (HgS) in a glass tube open at both 
 ends; clamp the tube in an inclined position so as to permit a free draught 
 of air through it. Heat the tube strongly with the Bunsen flame. Results? 
 Identify the fumes ( ?). Equation ? 
 
 Experiment III. Properties of the Salts of Mercury. 
 
 Examine the different salts of mercury (end shelf), noting the obvious 
 physical properties. Heat a small quantity of a mercury salt in a test 
 tube. Result? This property of volatilizing unchanged, giving sub- 
 limates of the same compound, is characteristic of many mercury com- 
 pounds. Give the name, formula and color of each salt. How many 
 series of mercury salts do you find ? Tabulate the foregoing data. 
 
 Give the formula and use of each of the following substances: calo- 
 mel, corrosive sublimate, vermillion. 
 
 Experiment IV. Formation of Mercurous Nitrate from Dilute Nitric 
 Acid and Mercury. 
 
 By means of a glass pipette take two or three drops of mercury from 
 the supply-bottle, place the mercury in a test tube or beaker, and add 
 15 cm. 3 of dilute nitric acid (i to i). Allow the action to continue for 
 about an hour (heating gently will hasten the action). An excess of 
 mercury must always be present. Continued stirring will cause crystal- 
 lization of the mercurous nitrate which is formed in the solution under 
 the above conditions. Dissolve the crystals in distilled water which has 
 been made slightly acid by the addition of a few drops of nitric acid. 
 Filter the solution, and reserve the filtrate for use in those experiments 
 in which mercurous nitrate is demanded. Equation? What is the 
 valence of mercury in mercurous nitrate ? 
 
 Experiment V. Formation of Mercuric Nitrate from Concentrated 
 Nitric Acid and Mercury. 
 
 Boil 10 cm. 3 of mercurous nitrate or two or three globules of mercury 
 with an excess of strong nitric acid. Continue the boiling until all 
 of the mercury has disappeared, evaporate nearly to dryness upon a 
 steam bath, add a few drops of nitric acid, and continue the evaporating 
 until the reddish-brown fumes cease to be liberated. The residue is 
 mercuric nitrate. Prepare a solution of the salt according to the direc- 
 tions given in the preceding experiment. Reserve the solution for 
 experiments. 
 
324 EXPERIMENTAL CHEMISTRY. 
 
 Experiment VI. Precipitation of Mercurous Iodide and Mercuric 
 Iodide. 
 
 (a) Add a solution of potassium iodide drop by drop to a dilute 
 solution of mercurous nitrate. Result ? Equation ? What is the effect 
 when an excess of the precipitant is added ? 
 
 (b) Repeat (a) using mercuric nitrate instead of mercurous nitrate. 
 Observe the precipitation of the yellow variety of mercuric iodide, which 
 in a few minutes change into the red variety. The conversion into the 
 red form is greatly hastened by light. Add an excess of the solution of 
 potassium iodide. Result ? What complex ion is formed ? Equation ? 
 
 Note. When the red iodide of mercury is heated above 126 it is 
 converted into the yellow crystalline variety of mercuric iodide which if 
 kept in the cold, changes slowly into the red. If the yellow variety is 
 scratched or rubbed, it is converted at once into the red crystalline form. 
 The red variety is stable at temperatures below 126, and the yellow, 
 above 126. This is the transition temperature which separates the two 
 regions of stability. Substances like sulphur and mercuric iodide which 
 can change in two directions-i.e., have two regions of stability-are said 
 to possess the property of " enantiotropy."* 
 
 Experiment VII. Analytical Reactions of Mercury Salts. 
 
 (a) Heat a small quantity of a salt of mercury with an equal amount 
 of anhydrous sodium carbonate in a dry test tube. Result ? 
 
 (b) Recall or repeat Exp. XI (d) " Electrolysis and Electrical 
 Equivalents." 
 
 Note. To separate portions of a dilute solution of mercurous nitrate 
 add the following reagents. Repeat each experiment using a dilute 
 solution of any mercuric salt. Compare and tabulate the results. 
 
 (c) Test with litmus paper ( ?). Explain. 
 
 (d) A solution of sodium hydroxide (?). Equations? 
 
 (e) Ammonium hydroxide (?). Equations? 
 
 (/) Dilute hydrochloric acid (?). If a precipitate is formed, filter 
 and treat the precipitate upon the filter paper with ammonium hydroxide. 
 Equations ? 
 
 Note. The compound formed by the action of the ammonium hy- 
 droxide is mercurous chloramide, NH 2 .Hg 2 .Cl. A compound formed 
 by the action of ammonium hydroxide on mercuric chloride is known 
 as mercuric chloramide, NH 2 .Hg.Cl. 
 
 (g) Hydrogen sulphide ( ?). If a precipitate is formed, ascertain its 
 solubility in strong nitric acid. Equation? 
 
 (h) Stannous chloride (?). Use an excess of the reagent and warm 
 gently (?). Equations? 
 
 * Ostwald Principles of Inorganic Chemistry. 
 
MERCURY. 325 
 
 (i) Compare the properties of the monomercurion, Hg', and the dimer- 
 curion, Hg". 
 
 (/) State how a solution of a mercurous compound can be distin- 
 guished from a solution of a mercuric salt. 
 
 How could you separate the metals from a solution containing the 
 salts of zinc, cadmium and mercury? Make a brief coherent report 
 of the method of procedure. 
 
 PROBLEMS. 
 
 1. The atomic weight of mercury is 200.0. If the vapor density of 
 mercury is 100, how many atoms does the molecule contain ? 
 
 2. 10 grams of zinc will liberate from sulphuric acid how many cm. 3 
 of hydrogen at o C. and 760 mm. ? At 20 C., and 730 mm. ? 
 
 3. Calculate the percentage composition of cadmium sulphide. 
 
CHAPTER XXVIII. 
 THE ELEMENTS OF GROUP III. 
 
 Family M. Family m. 
 
 (Scandium, Sc. 44. ) Boron, B. n.o 
 
 (Yttrium, Y. 89.0) Aluminum, Al. 27.1 
 
 (Lanthanum, La. 138.9) (Gallium, Ga. 70.0) 
 
 (Ytterbium, Yb. 173.0) (Indium, In. 115. ) 
 
 Thallium, Tl. 204.1 
 
 Scandium, yttrium, lanthanum and ytterbium, like certain other 
 elements of the following groups, are generally known as rare earth 
 metals. Cerium (Ce.; At. Wt.i4o.2), praseodymium (Pr.; At. VVt. 140.5), 
 neodymium (Nd.; At. Wt. 143.6), samarium (Sa.; At. Wt. 150.3), gado- 
 linium (Gd.; At. Wt. 156), and erbium (Er.; At. Wt. 166) are usually in- 
 cluded amongst the metals of the rare earths. They have been given 
 this generic name because they are found only in small amounts in a few 
 rare earthy minerals, such as orthite, euxenite, cerite, gadolinite and 
 monazite. Very few of these metals have been isolated in pure elementary 
 form, but many of their salts have been prepared. The latter, however, 
 are so similar in behavior that separation is exceedingly difficult. 
 
 The study and isolation of the elements of this group is attended 
 by much difficulty. This is well illustrated by the history of the isolation 
 of the compounds of the elements of praseodymium and neodymium. 
 In 1839, Mosander prepared from the mineral yttria, certain oxides, 
 among which there was one which he regarded as a compound of a new 
 elementary substance, to which the name didymium was eventually 
 given. Nearly sixty years elapsed before the elementary character of 
 didymium was questioned. The comparatively recent researches of 
 Auer von Welsbach have shown, however, that didymium is a mixture 
 of the two elements, praseodymium and neodymium. The salts of 
 praseodymium are light green while the neodymium salts possess a 
 rose-violet color. 
 
 The properties of these elements are similar to those of aluminum, 
 showing the same gradations as w r ere found in preceding groups, with 
 increase of atomic weights; that is, the free metals possessing the greater 
 combining weights are the more readily oxidized and their respective 
 bases are stronger. In their compounds the metals are usually trivalent. 
 The oxides and hydroxides are insoluble in water, and the chlorides, 
 nitrates and sulphates are soluble. The nitrates are decomposed by 
 heat. "Many of these elements give a very complicated spectrum on 
 allowing the electric spark to pass between carbon points moistened with 
 solutions of their salts. Since, under given conditions, each element 
 
THE ELEMENTS OF GROUP III. 
 
 327 
 
 possesses a perfectly definite spectrum, it can be seen whether the spec- 
 trum changes by partial separations. Where this is the case, we are 
 certainly dealing with a mixture. The higher members also exhibit 
 absorption spectra, some of them also emission spectra." 
 
 Several of the metals of this group have atomic weights so near to 
 lanthanum that it seems impossible at present to accommodate them 
 in the periodic table. Ostwald.* in discussing the probable position of 
 lanthanum in the table, says: " This signifies that there exist at this point 
 not one element, but a number of elements which are all very close to one 
 another, and have therefore an almost equal claim to this position. This 
 is an occurrence of numerous small planetary bodies at a point of the 
 solar system where, by analogy, one would have expected a large planet." 
 
 Scandium, which was discovered in 1879 by Nilson and Cleve, is of 
 particular interest to us because its existence was suggested in 1869 by 
 Mendeleeff who, from considerations based upon the periodic table, 
 predicted its atomic weight and many of its physical and chemical 
 properties. Mendeleeff had given the name of eka-boron to the element. 
 
 Eka-boron (Predicted 1869). 
 
 Atomic weight, about 44. 
 
 Oxide, Eb 2 O 3 , soluble in acids, 
 analogous to A1 2 O 3 , but more 
 basic; insoluble in alkalies. 
 
 Scandium (Discovered 1879). 
 
 Atomic weight, 44.1. 
 
 Oxide, Sc 2 O 3 , soluble in strong 
 acids, analogous to A1 2 O 3 , but 
 much more basic; insoluble in 
 alkalies. 
 
 Salts colorless and give gelatin- Salts colorless and give gelatinous 
 ous precipitates with NaOH or precipitates with NaOH or 
 
 Na 2 CO 3 . 
 
 Sulphate, Sc 2 (SO 4 ) 3 , forms a 
 double salt with K 2 SO 4 , which is 
 not isomorphous with the alums. 
 
 Gallium (Discovered 1875). 
 Atomic weight, 69.9. 
 Melting-point, 30.2. 
 Specific gravity, 5.93. 
 
 Slightly oxidized at red heat. 
 Decomposes water at high tem- 
 perature. 
 
 Gallium oxide, Ga 2 O 3 . 
 Gallium chloride, Ga 2 Cl 6 . 
 Gallium sulphate, Ga 2 (SO 4 ) 3 . 
 Forms a well-defined alum. 
 
 Sulphate, Eb 2 (SO 4 ) 3 , will form a 
 double salt with K 2 SO 4 , not iso- 
 morphous with the alums. 
 
 Eka-aluminum (Predicted 1871). 
 Atomic weight, about 69. 
 Will have a low melting-point. 
 Specific gravity, about 5.9. 
 
 Will not be acted upon by the air. 
 Will decompose water at red heat. 
 
 Will give an oxide, E1 2 O 3 . 
 Will give a chloride, E1 2 O 6 . 
 Will give a sulphate, E1 2 (SO 4 ) 3 . 
 Will form a potassium alum, etc. 
 
 *Principles of Inorganic Chemistry. 
 
328 EXPERIMENTAL CHEMISTRY. 
 
 With the exception of boron, aluminum and thallium, the members 
 of Family m. are among the rarest of the elements. Gallium, indium 
 and thallium were discovered by means of the spectroscope; gallium 
 takes its name from the country (France) in which it was discovered by 
 Lecoq. de Boisbaudrau, in 1875; indium, by Reich and Richter in 1863, 
 received its name on account of two characteristic lines in the indigo- 
 blue part of the spectrum; and thallium, by Crookes in 1861, owes its 
 name to the fact that there is a prominent green line in its spectrum 
 (Gk. OaXXSf, a green twig). These metals occur in exceedingly small 
 quantities as impurities in zinc blends. They are moderately heavy, 
 similar to one another in most properties, and possess specific gravities 
 and melting points which vary in the order of their atomic weights. 
 
 Gallium is a lustrous gray metal which possesses the extremely low 
 fusing-point of 30.2 C. It is oxidized superficially when heated to high 
 temperatures in the air, and is acted upon by strong acids and strong 
 basic hydroxides. Indium is a soft white metal which melts at 176. 
 It burns with a violet flame when heated in the air, and is acted upon by 
 strong acids. Thallium is a soft grayish-white metal which oxidizes 
 at ordinary temperature in moist air. It melts at 290. Sulphuric acid 
 and nitric acid act upon thallium, but hydrochloric acid acts upon it 
 only superficially as the insoluble thallous chloride which forms upon the 
 surface protects the metal. 
 
 The hydroxides of these three elements are weak bases. Gallium 
 hydroxide and indium hydroxide may also interact with strong basic 
 hydroxides, but thallium hydroxide shows no such acid properties. 
 Gallium functions both as a bivalent and a trivalent element; indium, 
 as a univalent, bivalent and trivalent element, and thallium, as a univalent 
 and a trivalent element. 
 
 Boron is the only non-metal of the group, all the others exhibit well 
 marked basic properties. The oxide, B 2 O 3 , of boron is acidic. (See 
 Chapter XXII). 
 
 Aluminum does not occur native, but its oxides and silicates are found 
 widely distributed. It is prepared by the electrolysis of oxide of alumi- 
 num dissolved in cryolite. Aluminum is a silver-white metal, very 
 ductile and malleable, a good conductor of electricity, and melts at 657. 
 It does not tarnish and is practically without action upon water. It is 
 scarcely acted upon by nitric acid, but readily dissolves in hydrochloric 
 acid and in strong solutions of the caustic alkalies with liberation of 
 hydrogen. 
 
 Aluminum functions as a trivalent element in its compounds. It is a 
 highly electro-positive element, and in consequence of its great affinity 
 for oxygen, aluminum displaces all metals save magnesium from their 
 oxides. The extreme readiness with which aluminum is able to effect 
 such reduction when the reaction is once started by heat and the ex- 
 ceedingly high temperature which is reached by the action have led to 
 some very useful applications. The heat of reaction between powdered 
 aluminum and ferric oxide has found useful application in the "Gold- 
 
THE ELEMENTS OF GROUP III. 329 
 
 schmidt process" for welding iron and steel. The very high temperature, 
 25oo-3ooo, produced by the reaction is sufficient to melt both the 
 iron and the oxide of aluminum. As the products of the action are not 
 miscible, they separate into two layers. The sulphides are reduced 
 just as readily. This furnishes a simple method for the preparation of 
 pure specimens of the metals whose oxides and sulphides are reduced 
 with difficulty. Goldschmidt, the inventor, has given to these processes 
 the general name of " ahiminothermy" Mixtures of the metallic oxides 
 and granulated aluminum have been placed on the market under the 
 name of "thermit" 
 
 The hydroxide of aluminum is feebly acidic as well as basic, and 
 therefore forms two classes of compounds of the types Na 3 AlO 3 and 
 A1C1 3 . The general name of aluminates is given to the first class. Both 
 series of salts are hydrolyzed by water, which is to be expected. 
 
 Aluminum has a wide range of uses, although the marked influence 
 of traces of impurities has limited its application more than was anticipated 
 originally. It is used for making aluminum bronze, magnalium, flash- 
 light powders, paint, cables for conducting electricity, foil, various 
 ornamental articles and cooking utensils. 
 
 The compounds of aluminum are of equal importance. Ruby, corun- 
 dum and sapphire are nearly pure oxides of aluminum. Emery is a 
 mixture of corundum and iron. Clay, an impure silicate of aluminum, 
 is used in the manufacture of bricks and earthenware. Kaolin, nearly 
 pure aluminum silicate, is used in the manufacture of porcelain and 
 china. Hydraulic cement (hardens under water) is made by heating a 
 mixture of limestone and clay to incipient fusion. 
 
 BORON, B. 
 
 At. Wt. n.o Sp. Gr. 2.5. 
 (See Chapter XXII). 
 
 ALUMINUM, Al. 
 At. Wt. 27.1 Sp. Gr. 2.58. 
 Experiment I. Properties of Aluminum. 
 
 (a) Examine specimens of aluminum in the forms of ingot, wire and 
 powder. Introduce a piece of the wire into the Bunsen flame in order 
 to determine whether aluminum is a conductor of heat (?). Does the 
 wire melt in the flame (use tongs)? Try the blast- lamp (?). What is 
 the melting point of the metal? Does it tarnish in the air? 
 
 (b) Solubility in acids. Try the action of the powdered or granulated 
 metal on the following reagents: hydrochloric acid (?), nitric acid ( ?), 
 and sulphuric acid (?). Account for results. Test the gas, if any is 
 liberated (?). Equations? 
 
 (c) Solubility in the caustic alkalies. Add about 5 cm. 3 of a con- 
 
330 EXPERIMENTAL CHEMISTRY. 
 
 centrated solution of sodium hydroxide or potassium hydroxide to i to 2 
 cm. 3 of granulated aluminum, and warm gently (?). Identify the gas 
 evolved (?). Add more of the hydroxide if necessary. The solution 
 contains sodium aluminate, Na 3 AlO 3 . Ionic equation ? Does the acidic 
 or basic properties of the aluminum hydroxide predominate in this 
 reaction? Your reasons? To prove that aluminum has gone into 
 solution, neutralize the product of the above reaction carefully with 
 dilute hydrochloric acid. Result ? Equation ? Are the properties of the 
 aluminum revealed by this reaction ? Explain. 
 
 (d) Tabulate the properties of aluminum. Should commercial 
 "lyes" (alkalies) be warmed in cooking utensils made of aluminum? 
 Why? 
 
 (e) Enumerate some of the uses of aluminum. 
 
 Experiment II. Reduction of Metallic Oxides by Use of Aluminum. 
 The Goldschmidt Process. "Thermit." 
 
 Thoroughly mix small quantities of iron oxide and granulated alumi- 
 num. Place the mixture on an iron plate or a piece of tile, and ignite it 
 by means of a piece of magnesium ribbon ( ?). Examine the fused mass 
 for globules of iron. If the above mixture does not ignite readily, procure 
 a small quantity of commercial mixture known as "thermit iron" 
 (iron oxide and aluminum) and ignite it after receiving instructions from 
 the assistant. Account for the results. Equation ? 
 
 Experiment III. Salts of Aluminum. 
 
 Examine the various aluminum salts found in the laboratory. Record 
 the name, formula and color of each salt. What is ultramarine, its 
 formula, color and use? 
 
 Experiment IV. Preparation of the Compound Salt, Potassium- 
 Aluminum Sulphate. Alums. 
 
 (a) Prepare hot saturated solutions of potassium sulphate and alumi- 
 num sulphate 25 cm. 3 of each solution will be sufficient. Mix the 
 two solutions in a small beaker or crystallizing dish so that the 
 resulting solution contains the weights of the two salts in a proportion 
 approximating the ratio of their molecular weights (10 grams of 
 Al 2 (SO 4 ) 3 .i8H 2 O and 3 grams of K 2 SO 4 in 200 cm. 3 of distilled water. 
 Concentrate to 100 cm. 3 Cool.) The compound salt, potassium- 
 aluminum sulphate, K 2 SO 4 , A1 2 (SO 4 ) 3 . 24H 2 O, crystallizes on standing 
 (it may be necessary to allow it to stand for several days). What is the 
 form of the crystals ? 
 
 (b) Repeat (a), using ammonium sulphate instead of potassium sul- 
 phate. Result? Compare the forms of the crystals (?). 
 
 (c) What is an alum ? State some of the uses of alums. Notice the 
 composition of the various alums, a list of which is appended. All of 
 these alums crystallize in the same form and possess the same funda- 
 mental chemical poperties. 
 
THE ELEMENTS OF GROUP III. 331 
 
 Na2SO 4 .Al 2 (SO 4 ) 3 .24H 2 O, Sodium alum. 
 K 2 SO 4 .A1 2 (SO 4 ) 3 .24H 2 O, Potassium alum. 
 (NH 4 ) 2 SO 4 . A1 2 (SO 4 ) 3 .24H 2 O, Ammonium alum. 
 
 Alums are also formed by substituting iron, chromium and man- 
 ganese for aluminum. 
 
 K 2 SO 4 .Fe 2 (SO 4 ) 3 .24H 2 O, Iron alum. 
 K 2 SO 4 . Cr 2 (SO 4 ) 3 . 2 4 H 2 O, " Chrome " alum. 
 K 2 SO 4 . Mn 2 (SO 4 ) 3 . 24H 2 O, Manganese alum. 
 
 Experiment V. Analytical Reactions of Aluminum Salts. 
 
 (a) Heat a small quantity of an aluminum salt (use alum) on a piece 
 of charcoal in the oxidizing flame (?), moisten with a few drops of a 
 cobalt nitrate solution, and heat again. Account for the blue color of 
 the product. 
 
 (b) Test a solution of the chloride or sulphate of aluminum with litmus 
 paper (?). Explain. 
 
 (c) To 10 cm. 3 of an aluminum sulphate solution add a small quantity 
 of a solution of sodium hydroxide. Filter off the precipitate, and suspend 
 it 'in water. Divide the mixture into two portions. To one portion 
 add hydrochloric acid (?). Equations? To the second portion add a 
 slight excess of the sodium hydroxide solution (?). Now add an excess 
 of a solution of ammonium chloride ( ?). Account for the re-precipi- 
 tation of aluminum hydroxide. Equations? 
 
 (d) Add a sodium carbonate solution to a solution of aluminum 
 sulphate (?). Filter off the precipitate, and wash it until free from 
 sodium carbonate. Test the precipitate with hydrochloric acid. Is it 
 soluble and is carbon dioxide evolved. Explain. Equations? 
 
 (e) Add ammonium sulphide to a solution of an aluminum salt. 
 Result? Filter, and wash the precipitate until it is free from the pre- 
 cipitant, i.e., is odorless. Remove the precipitate to a test tube, and 
 add hydrochloric acid (?). Is hydrogen sulphide evolved? Was the 
 precipitate a sulphide? 
 
 The hydroxide and not the sulphide was formed in the first reaction 
 as the sulphide cannot exist in the presence of water. Explain. Express 
 the precipitation of aluminum hydroxide by ammonium sulphide by 
 ionic equations. 
 
 (/) Mordants, (i) To i cm. 3 of a cochineal solution add 5 or 6 cm. 3 
 of a solution of aluminum sulphate, then add ammonium hydroxide. 
 Shake vigorously and filter. Locate the coloring matter (?). Does 
 water remove it easily? 
 
 (2) Immerse a strip of white muslin in a strong solution of aluminum 
 sulphate. When the muslin becomes saturated, transfer it to a hot 
 solution of cochineal which has been made strongly alkaline with am- 
 monium hydroxide. What are mordants? State the purpose of the 
 use of the ammonium hydroxide in conjunction with the aluminum 
 sulphate. 
 
332 EXPERIMENTAL CHEMISTRY. 
 
 THALLIUM, Tl. 
 
 At. Wt. 214.1 Sp. Gr. U.S. 
 Experiment I. Properties of Thallium. 
 
 Examine a specimen of the metal (?). It tarnishes on exposure to 
 air with the formation of black thallous oxide. Has the specimen under- 
 gone oxidation ? 
 
 In the univalent condition its properties suggest those of sodium and 
 silver, while trivalent thallium in its compounds resembles aluminum; 
 for example, thallic salts are hydrolyzed by water. 
 
 Experiment II. Salts of Thallium. 
 
 Examine those compounds of the metal which are found on the end 
 shelf. Observe the name, formula and color of each salt. How many 
 well-defined series of thallium compounds do you find? What is the 
 valence of thallium in each series? Tabulate data. 
 
 PROBLEMS. 
 
 i. Calculate the amount of alumnia, A1 2 O 3 , which can be prepared 
 from 1000 grams of bauxite, A1 2 O(OH) 4 . 
 
 A1 2 O(OH) 4 + Na^Og 2NaAlO 2 + CO 2 + 2H 2 O, 
 2NaAlO 2 + CO 2 + 3H 2 O -> N 2 CO 3 + 2A1(OH) 3 . 
 2 A1(OH) 3 -> A1 2 O 3 + 3H 2 O. 
 
CHAPTER XXIX. 
 THE ELEMENTS OF GROUP IV. 
 
 Family M. 
 (Titanium, Ti. 
 (Zirconium, Zr. 
 (Cerium, 
 (Thorium, 
 
 48.1 
 90.6 
 
 Ce. 140.25 
 Th. 232.5 
 
 Family m. 
 Carbon, C. 
 
 Silicon, Si. 
 
 (Germanium, Gr. 
 Tin, Sn. 
 
 Lead, Pb. 
 
 12 .OO 
 
 28.4 
 
 72.5 
 II9.O 
 206.9 
 
 The four rare elements, titanium, zirconium, cerium and thorium, 
 compose Family M of Group IV. 
 
 Titanium as the constituent of several rare minerals occurs chiefly in 
 the form of the crystalline titanium oxide, of which there are several 
 native varieties, known as rutile, anatase and brookite. It is exceedingly 
 difficult to prepare the metal in the pure form, owing to the fact that it 
 combines readily with nitrogen, forming a nitride. Zirconium is found 
 in the rare silicate, ZrSiO 4 , in the mineral zircon. This element resembles 
 boron in that it occurs in several allotropic varieties: an amorphous form, 
 a black powder, a graphitic form, and the crystalline form, hard steel- 
 gray lamella. The amorphous form burns in the air when heated 
 gently, but the crystalline varieties require the high temperature of the 
 oxyhydrogen flame for its ignition. Zirconium dioxide (zirconia), 
 ZrO 2 , which emits a white and brilliant light when heated, is used as a 
 constituent of the mantles of the Welsbach lights. Cerium occurs chiefly 
 in the two rare minerals, cerite and orthite. The metal as obtained by 
 electrolysis is steel-gray, lustrous and malleable. It is stable in the 
 air at ordinary temperatures, but possesses the peculiar property of 
 undergoing oxidation with incandescence when heated. In virtue of 
 this latter property, cerium has recently become of much technical 
 importance in the manufacture of the Welsbach mantle. It has been 
 suggested* that " the effectiveness of cerium dioxide, CeO 2 , as a con- 
 stituent of the Welsbach mantle in intensifying the light emitted may 
 depend upon the facility with which it is reduced and reoxidized, thereby 
 acting as a catalyzer to cause rapid combustion and high temperature 
 in the flame." Thorium, which until comparatively recently was obtained 
 almost solely from thorite (a silicate), is now obtained from the mineral 
 monazite which is found in abundance in North Carolina. Its metallic 
 properties are very similar to those of cerium. The dioxide of thorium 
 is the chief constituent of the Welsbach mantles, which are usually pre- 
 pared by moistening a cotton web with a solution containing thorium 
 
 * Gooch and Walker Outlines of Inorganic Chemistry. 
 333 
 
334 EXPERIMENTAL CHEMISTRY. 
 
 nitrate and one per cent, of cerium nitrate, then drying and heating it. 
 The web (organic material) burns and the thoria and ceria remain as a 
 white coherent frame-work. The mantle is made ready for use by 
 heating it in a Bunsen fla.me produced under pressure. The material 
 of the mantle contracts appreciably as the result of this last operation. 
 
 Although the exact role of the ceria is not known, yet it is a fact that 
 pure thoria makes a poorly luminous mantle. Ostwald* says: "The 
 cause of this influence (referring to the luminous properties of the mantle 
 when small quantities of certain other substances are mixed with the 
 thoria) has not yet been established quite free from doubt, but the most 
 probable view is that the addition (of the ceria) effects a catalytic acceler- 
 ation of the combustion of the mixture of coal-gas and air in direct contact 
 with the skeleton of thoria. Besides this, the optical properties of thoria 
 appear to be of importance." 
 
 The compounds of thorium are radio-active; i.e., they possess the 
 power of continually emitting certain "influences," sometimes called 
 the "Becquerel rays" and "emanations" which possess the property 
 of effecting a change in the electrical properties of the air. Certain of 
 these "influences" are able to penetrate solid substances and affect a 
 photographic plate. The student is referred to reference works on this 
 subject of radio activity, f 
 
 All the metals of this family are quadrivalent. 
 
 The elements of Family m. have been divided into two series (see 
 Chapter XX); a primary series embracing carbon and silicon, and a 
 secondary series composed of the elements, germanium, tin and lead. 
 The primary series has been considered previously. 
 
 The elements of the secondary series are silver-w r hite, lustrous, mal- 
 leable metals which are practically unaffected by air or water. The 
 metals are fusible and volatile at high temperatures. They possess 
 high specific gravities which increase as the atomic weights increase. 
 
 The rare element germanium (Mendeleeff's eka-silicon) forms a 
 sort of link between the elements of the primary series and those of 
 the secondary series. It is both metallic and non-metallic; its oxide 
 combines with acids, but it also unites with alkaline hydroxides. Both 
 tin and lead which resemble each other, especially in their physical 
 properties, show similar basic and acidic relations in their respective 
 compounds; i.e., their oxides and hydroxides combine with strong acids 
 to form salts in which tin and lead are positive ions, while with the 
 caustic alkalies they form compounds, known respectively as stannates 
 of the type, NaSnO 2 , and plumbites of the type, Na-jPbO-j, in which the 
 metals are in the negative ions. Each element of the series forms com- 
 pounds of the types stannous chloride (SnCl 2 ) and stannic thloride (SnCl 4 ), 
 in which the metals function, respectively, as bivalent and quadrivalent 
 
 * Prin. of Inorg. Chem. 
 
 t Rutherford Radioactivity; Arrhenius Theories of Chemistry; Fournier The 
 Electron Theory. 
 
THE ELEMENTS OF GROUP IV. 335 
 
 elements. Lead, however, manifests a greater tendency to react at the 
 state of lower valency. This is not an exception, but rather a rule that 
 the heavier metals of a group tend to react at the lower valency. 
 
 The compounds of these three elements, many of which are insoluble, 
 are but slightly hydrolyzed by water and generally stable. 
 
 The chief ore of tin is tin-stone, or cassiterite, SnO 2 , the greater portion 
 of which (fully 80 per cent.) is supplied by Cornwall and the East Indies. 
 The metallurgy of tin consists of four processes, namely, (i) crushing 
 or pulverizing the ore, (2) calcining, (3) washing, (4) smelting or re- 
 ducing. After the ore has been crushed finely and washed to free it 
 from earthy matter, it is calcined in a reverberatory furnace. The 
 latter operation is for the purpose of oxidizing the sulphides of iron and 
 copper and to drive off the arsenic. The sulphur and arsenic are led 
 into condensing flues where the arsenic deposits are collected. The 
 calcined ore is now washed to eliminate the oxide of iron and the sulphate 
 of copper, and then reduced with powdered anthracite coal in a rever- 
 beratory furnace. The tin obtained by the foregoing process is further 
 purified by remelting at a gentle heat, the pure and more readily fusible 
 tin being allowed to flow away from the residue and alloys of other 
 metals. 
 
 The crystalline character of the metal may be observed by pouring 
 warm dilute aqua regia over the surface of a piece of block-tin or a sheet 
 of tinned iron. The surface of the metal after such treatment exhibits 
 a beautiful crystalline appearance. Again, when a bar of tin is bent 
 a peculiar crackling sound (tin cry) is produced. It has been observed 
 also that the metal becomes perceptibly hot at the place of flexure. It 
 has been suggested that the cause of these phenomena may be due to the 
 friction of the crystals upon one another. When ordinary tin, which 
 has a specific gravity of 7.3. is exposed to the prolonged influence of low 
 temperature, it changes to a gray pulverulent variety of specific gravity 
 5.8. The transition temperature is 20, and ordinary tin is in a meta- 
 stable condition below this temperature. Tin is not tarnished by air 
 or water at ordinary temperatures, but when heated above the melting- 
 point it burns with a brilliant white light, forming white clouds of stannic, 
 SnO 2 . At red heat tin decomposes steam with a liberation of hydrogen. 
 Hydrochloric acid and tin interact with the evolution of hydrogen and 
 the formation of stannous chloride; when strong sulphuric acid is heated 
 w r ith tin, stannous sulphates and sulphur dioxide are formed; cold 
 dilute nitric acid yields stannous nitrate and ammonia, while concen- 
 trated nitric acid gives stannic nitrate, which is hydrolyzed by water with 
 the formation of metastannic acid (H 2 SnO 3 ) 5 . The caustic alkalies 
 attack tin, giving hydrogen and a metastannate, such as Na 2 SnO 3 . Tin 
 forms alloys with lead, copper, antimony, bismuth and mercury. Among 
 the more familiar of these alloys are bronze, sojt solder (50 per cent, 
 lead), pewter (25 per cent, lead), Britannia metal (10 per cent, antimony 
 and a small quantity of copper), and fusible alloys (bismuth, tin, cadmium 
 and sometimes lead). Tin is also used as a protective covering on other 
 
33 EXPERIMENTAL CHEMISTRY. 
 
 metals on account of the difficulty with which it is attacked by many 
 corroding substances. Tin-plate is made by dipping sheet iron into 
 molten tin. Vessels made of copper are also frequently covered. Or- 
 dinary brass pins are made of brass wire coated with tin. The metal 
 finds many other uses in the arts. 
 
 Lead is obtained almost wholly from the ore, galena PbS. The 
 metallurgical processes by which lead is obtained from its ores are 
 similar to those described for the reduction of sulphides. The metal 
 is refined by electrolytic methods. It is a grayish-white metal, soft and 
 tough. It melts at about 330, and vaporizes at 1700. Under suitable 
 conditions it crystallizes, the crystals having the octahedral form. When 
 warm, it may be formed into pipes by hydraulic pressure. It is a poor 
 conductor of electricity. Lead is oxidized but superficially by the air, 
 becoming covered with a film of a dark-colored oxide, which is probably 
 the suboxide Pb 2 O, the composition of the final covering being that of a 
 basic carbonate; when heated, the metal passes through several stages of 
 oxidation with the formation of no less than five distinct oxides. Pure 
 water does not act upon lead, but hard water covers it with a coating 
 composed largely of the sulphate and the carbonate. As these salts 
 are insoluble, they protect the metal and prevent contamination of the 
 water with poisonous lead compounds. Water holding air in solution 
 attacks lead, forming the slightly soluble hydroxide and the carbonate. 
 The latter is appreciably soluble in water containing carbon dioxide. 
 The use of lead pipes for conducting water may become a source of 
 danger, as rain water especially, owing to the presence of oxygen and 
 carbon dioxide, is likely to exert a solvent action upon the pipes. The 
 lead compounds taken continuously into the system in small quantities 
 act as a cumulative poison. Hydrochloric acid acts upon the metal 
 slowly with an evolution of hydrogen. Concentrated sulphuric acid 
 has but little affect upon it; dilute sulphuric acid slowly interacts with it, 
 forming the insoluble lead sulphate and sulphur dioxide. Nitric acid 
 attacks it, giving lead nitrate and oxides of nitrogen. The action of the 
 caustic alkalies upon lead has been mentioned previously. 
 
 The uses of lead are numerous. It forms useful alloys with tin, 
 antimony, bismuth, copper and zinc. Type-metal usually contains lead 
 and antimony and sometimes tin. On account of the resistance of lead 
 to the action of air and water and many other substances, it is employed 
 in making various kinds of vessels and lead pipes for carrying water. 
 Shot-metal is an alloy composed of about 99.5 per cent, of lead and 0.5 
 per cent, of arsenic. " In the process of making shot the lead is melted 
 in a cast-iron pan, and after the addition of a sufficient amount of lead 
 arsenide to form an alloy containing from 0.8 to 0.9 per cent, of arsenic, 
 the molten alloy is poured into a perforated iron basin at the top of a 
 high tower, and allowed to fall into a dilute solution of sodium sulphide. 
 The presence of arsenic makes the drops of molten alloy very fluid, so 
 that they assume a spherical form in their passage to the bottom of the 
 tower. The effect of the sodium sulphide is to coat the shot with a thin 
 
THE ELEMENTS OF GEOUP IV. 
 
 337 
 
 layer of lead sulphide which prevents superficial oxidation when the shot 
 are removed from the water. Shot are sometimes made by allowing the 
 molten metal to fall through an ascending air current or by pouring 
 it in a thin stream upon a rapidly revolving disk, when the centrifugal 
 action divides the metal into drops, which are thrown against a surround- 
 ing screen."* 
 
 The following table gives a general view of the physical and chemical 
 properties of the metals of this series: 
 
 Atomic weight, 
 Specific gravity, 
 Melting-point, 
 Volatilizes, 
 
 Oxides, 
 
 Chlorides, 
 
 Hydroxides, 
 Sulphides, 
 
 Germanium. 
 
 7 2 -5 
 
 5-47 
 900 
 
 Tin. 
 
 119.0 
 
 231 
 
 I S 00(?) 
 
 Lead. 
 206.9 
 
 n-37 
 330 
 1700 
 
 GeO; GeO 2 , SnO; SnO 2 , Pb 2 O; PbO; Pb 2 O 3 ; 
 
 PbO 2 ; Pb 3 O 4 . 
 
 GeCl 2 ;GeCl 4 , SnCl 2 ; SnCl 4 , PbCl 2 ; PbCl 4 . 
 
 Ge(OH).,; Ge(OH 4 , Sn(OH) 2 ,Sn(OH) 4 , Pb(OH) 2 . 
 GeS; GeS 2 SnS; SnS 2 , Pb 2 S; PbS. 
 
 GERMANIUM, GC. 
 
 At. Wt. 72.5 Sp. Gr. 7.3. 
 (See lecture notes and text-book.) 
 
 At. Wt. 119.0 Sp. Gr. 7.3. 
 Experiment I. Properties of Tin. 
 
 (a) Examine a specimen of this metal and record its most obvious 
 physical properties. What is its melting point ? Is it hard or soft ? 
 
 (b) (Quant.) Determine the specific heat of tin. Secure the directions 
 for procedure from the assistant. Calculate the atomic weight of tin, 
 using 0.054 as its specific heat. 
 
 (c) Does tin tarnish readily w r hen exposed to the action of the air? 
 Enumerate some of the uses of tin. Name its chief ores. 
 
 Experiment II. (Quant.) 
 Tin. 
 
 Determination of the Equivalent Weight of 
 
 In a crucible which has been previously desiccated and weighed, 
 place an accurately weighed quantity (about 0.5 gram.) of pure granulated 
 tin, and cover the metal with about 10 cm. 3 of concentrated nitric acid. 
 
 * Gooch and Walker. Outlines of Inorganic Chemistry. 
 
 22 
 
338 EXPERIMENTAL CHEMISTRY. 
 
 Carefully apply heat (use a hot iron plate) to the crucible. Identify the 
 reddish-brown fumes. When all of the tin has dissolved and the nitric 
 acid has been entirely expelled, place the crucible on a pipe-stem triangle 
 and heat it with a Bunsen burner. The white product is stannic oxide, 
 SnO 2 . Equation? Cool, and weigh the crucible and contents. Cal- 
 culate the chemical equivalent of tin. If the valence of tin in stannic 
 oxide is four, what is its atomic weight ? 
 
 Experiment III. Formation of Halides of Tin. 
 
 (a) Stannous chloride. Treat several small pieces of granulated tin 
 with 10 cm. 3 of concentrated hydrochloric acid in a test tube, and warm 
 gently to start the action. After the action has continued for some time, 
 pour off the liquid into another test tube and reserve it for use in those 
 experiments in which stannous chloride, SnCl. 2 , is required. Equation ? 
 
 (b) Stannic chloride. Add 3 cm. 3 of concentrated nitric acid and i 
 cm. 3 of hydrochloric acid to 5 cm* 3 of stannous chloride, and heat gently. 
 The solution contains stannic chloride, SnCl^. Dilute with 5 cm. 3 of 
 water. Equations ? 
 
 Experiment IV. Precipitation of Stannous Hydroxide. Sodium Stan- 
 nite. 
 
 To 5 cm. 3 of a dilute solution of stannous chloride add carefully a 
 sodium hyrdoxide solution until precipitation is complete. Equation ? 
 
 Divide the precipitate into two parts. To one portion add an excess 
 of the sodium hydroxide solution ( ?) and to the other portion add 
 hydrochloric acid (?). Equations? 
 
 Suggest a method for the preparation of stannous oxide, SnO. Equa- 
 tion? 
 
 Experiment V. Precipitation of Stannic Hydroxide (a Stannic Acid). 
 Sodium Stannate. 
 
 Repeat Experiment IV substituting stannic chloride for stannous 
 chloride. 
 
 What product is obtained by heating stannic hydroxide in a crucible ? 
 Equation ? 
 
 Epperiment VI. Analytical Reactions. 
 
 (a) Heat a small piece of tin on charcoal before the blow-pipe (?). 
 Allow a drop of cobalt nitrate to fall upon the incrustation, then heat 
 again (?). 
 
 (b) Place a piece of zinc or suspend a strip of it in a solution of 
 stannous chloride for ten to fifteen minutes. Result ? Equation ? 
 
 (c) Add i cm. 3 of a solution of mercuric chloride to 4 or 5 cm. 3 of 
 stannous chloride solution. Result? The solution now contains 
 stannic chloride, SnCl 4 . Equation? Warm the mixture and note the 
 changes ( ?). Equation ? Does the stannous chloride act as a reducing 
 or an oxidizing agent? Indicate all of the above changes by "ionic" 
 equations. Repeat the foregoing reaction, using stannic chloride instead 
 
THE ELEMENTS OF GROUP IV. 33 Q 
 
 of stannous chloride (?). How can you distinguish between stannous 
 and stannic compounds? 
 
 (d) Add a few drops of stannic chloride to 5 cm. 3 of dilute hydro- 
 chloric acid in a test tube. Now add several pieces of magnesium 
 ribbon to the acid solution. What is evolved? Is it an oxidizing or 
 reducing agent when in the nascent condition? When the magnesium 
 has dissolved, add a small quantity of the mercuric chloride solution. 
 Result ? Conclusions ? Equations ? 
 
 (e) Pass hydrogen sulphide into 5 cm. 3 of absolution of stannous chloride 
 containing i cm. 3 of dilute hydrochloric acid. Result? Equation? 
 Is the reaction easily reversible ? Give reasons for your answer. Filter 
 and wash the precipitate. Place the latter in an evaporating dish, add 
 10 cm. 3 of ammonium polysulphide and warm. Result? Equation? 
 Now add hydrochloric acid to the solution of ammonium sulphostannate, 
 (NH 4 ) 2 SnS 3 . Results? What gas was evolved? Equation? 
 
 Repeat the preceding reactions, using stannic chloride (?). 
 (/) There are how many ionic forms of tin ? What is the color of the 
 tin ions ? Give the valencies of the respective ions. 
 
 LEAD, Pb. 
 
 At. Wt. 206.9 Sp. Gr. 11.37. 
 Experiment I. Properties of Lead. 
 
 (a) File or scrape off the coating from a piece of lead. Is the metal 
 hard or soft? Color? Does lead tarnish (oxidize) readily in the air? 
 Try to mark on paper with lead ( ?). 
 
 (b) Solution tension. Dissolve about i gram of lead acetate, Pb(C 2 H 3 - 
 O 2 ) 2 , in 20 cm. 3 of water. Place a strip of sheet zinc or several pieces of 
 granulated zinc in the solution, and set aside for an hour. Result? 
 Ionic equation ? Remove the film of lead from the zinc, thoroughly 
 wash the former and reserve it for (c). 
 
 (c) Action of air and water on lead. Spread the finely divided lead 
 from (b) on a glass plate, moisten with a very little water, and expose to 
 the action of the air for an hour. Test the water with litmus papers ( ?). 
 Conclusions? The presence of lead in the water may be detected by 
 shaking a portion of the lead with a small volume of water and passing 
 hydrogen sulphide into the filtrate. A black precipitate (lead sulphide) 
 indicates the presence of lead. Equations ? 
 
 (d) Name the chief ore of lead. What are some of the uses of lead? 
 What is its melting point ? What is pewter ? 
 
 Experiment II. (L. T.) Precipitation of Lead from its Salts by Other 
 Metals. "Lead Tree." 
 
 Suspend a bar of zinc in a solution of lead acetate or nitrate. Allow 
 this to stand for several days. The lead will be gradually thrown out 
 of solution (why?) and deposited upon the zinc in arborescent forms, 
 
34 EXPERIMENTAL CHEMISTRY. 
 
 known as the "lead tree." Test the solution for the presence of zinc, 
 as follows: Add sufficient sulphuric acid to precipitate any lead remain- 
 ing in solution, filter, and to the filtrate add ammonia in excess and 
 ammonium sulphide; the zinc will be precipitated as sulphide. Ionic 
 equations ? 
 
 Experiment III. Salts of Lead. 
 
 Examine a number of the salts of lead. Give color and formula of 
 each. What is the valency of lead ? 
 
 Note. Lead forms five oxides; give the name and formula of each. 
 
 Experiment IV. Precipitation of Lead Hydroxide. Sodium Plumbate. 
 
 (a) To a solution of lead nitrate add at first slowly and then in excess 
 a sodium hydroxide solution. Describe all of the changes that occur. 
 Ionic equations? Does lead exhibit the properties of both a metal and 
 non-metal ? Give reasons for your answer. 
 
 (b) Repeat (a) using ammonium hydroxide. 
 
 Experiment V. Preparation of Lead Salts. 
 
 (a) Lead nitrate from lead and nitric acid. Treat about a gram of 
 lead in an evaporating dish with 20 cm. 3 of a mixture (i to 2) of nitric 
 acid and water. Place the dish on a wire gauze and heat gently until the 
 metal dissolves. Note the accompanying phenomena. Set the solu- 
 tion aside to crystallize by spontaneous evaporation. Result? Equa- 
 tion? 
 
 (b) Lead nitrate from lead monoxide (litharge) and nitric acid. Dis- 
 solve 3 to 5 grams of lead monoxide in dilute nitric acid, filter, evaporate 
 the filtrate to the crystallizing point, and cool. Result ? Equation ? 
 
 (c) Lead nitrate and lead dioxide from red-lead oxide (minium) and 
 nitric acid. Treat about a gram of minium with dilute nitric acid, 
 warm gently, and when the red color of the minium has changed to a 
 brown, dilute with water and filter. Wash the residue ( ?) and ignite 
 it. Which of the oxides of lead does it resemble most in appearance? 
 Prove the presence of lead in the filtrate by adding slowly an excess of a 
 solution of sodium hydroxide (?). Equations? This behavior of 
 minium suggests what theory, as to its constitution ? 
 
 (d) Lead acetate from lead oxide and acetic acid. To 5 grams of lead 
 monoxide add 10 cm. 3 of acetic acid, then warm gently to increase the 
 speed of the reaction. If the solution is not clear, filter, and evaporate 
 to the crystallizing point, exercising care to avoid charring the salt. 
 Equations? What is "sugar of lead?" 
 
 Experiment VI. Precipitation of Lead Carbonate. Basic Lead Car- 
 bonate ; its Decomposition by Heating. 
 
 (a) When a solution of ammonium carbonate is added to a solution of 
 lead nitrate, the normal carbonate is formed. (Equations ?) If any 
 other alkaline carbonate is used, a basic lead carbonate is formed. 
 
THE ELEMENTS OF GROUP IV. 341 
 
 (b) To 10 cm. 3 of a solution of lead nitrate add an excess of sodium 
 carbonate solution (?). Filter, wash and dry the precipitate, separate 
 it from the filter paper, and ignite the white powder in a porcelain crucible. 
 Avoid heating the crucible above dull redness. Observe the color of the 
 residue (?) when hot and cold. Equations? What is "white lead"? 
 Its use? 
 
 Experiment VII. Chemical Principles Involved in the Manufacture 
 of " White Lead." 
 
 Add a little lead monoxide to 10 cm. 3 of a solution of lead acetate; 
 thoroughly shake the mixture; pass carbon dioxide (generator on end 
 shelf) through it until it is white; filter. The precipitate is used as a 
 pigment under the name of "white lead." Equations? What is the 
 effect of hydrogen sulphide upon white lead ? Equation ? 
 
 Experiment VIII. Analytical Reactions. 
 
 (a) Heat a small piece of lead on charcoal in the oxidizing flame 
 (blow-pipe). Result? Equation? 
 
 (b) Test a solution of lead nitrate with litmus paper (?). Interpret 
 the result. 
 
 (c) Halides of lead. To 2 cm. 3 of a solution of lead nitrate add dilute 
 hydrochloric acid until precipitation is complete? Filter, wash the 
 precipitate, and heat it with the smallest volume of water which will 
 dissolve it. Cool the solution. Result ? Equations ? 
 
 In what two respects does lead chloride differ from silver chloride? 
 Name the insoluble chlorides. 
 
 Treat a cm. 3 of lead nitrate solution with a solution of potassium or 
 sodium bromide. Results ? Equation ? Is the precipitate soluble in 
 hot water ? Compare its solubility with that of lead chloride. 
 
 Add a solution of potassium iodide to 2 cm. 3 of lead nitrate solution. 
 Results ? Equation ? 
 
 How does the solubility of the precipitate compare with the solubility 
 of lead bromide? 
 
 (d) To 2 cm. 3 of lead nitrate solution add dilute sulphuric acid. Re- 
 sult? Equation. What sulphates are insoluble? Give the formula of 
 each. 
 
 (e) Allow hydrogen sulphide (generator on end shelf) to bubble 
 slowly through 5 cm. 3 of a solution of lead nitrate to which has been 
 added four or five drops of dilute nitric acid. Result? Equation? 
 Is the action reversed easily ? Give reasons for your answer. 
 
 (/) Add a dilute solution of potassium chromate to 3 cm. 3 of lead 
 nitrate solution. Result? Equation? Is the precipitate soluble in a 
 strong solution of sodium hydroxide? Equation? What is "chrome 
 yellow" ? Its uses? 
 
 (g) What is the color of the lead ion ? Does it ever form a part of the 
 anion ? If so, give an example. 
 
342 EXPERIMENTAL CHEMISTRY. 
 
 PROBLEMS. 
 
 1. Calculate the weight of tin in i kg. of cassiterite, SnO 2 . 
 
 2. What per cent, of lead is contained in galena, PbS ? 
 
 3. If 50 grams of tin yields 63.55 grams of stannic oxide, what is 
 the atomic weight of tin? 
 
 4. An analysis showed that i gram of lead monoxide contained 
 0.0717 gram of oxygen. Calculate the atomic weight of lead. 
 
 5. The per cent, of lead in lead chloride is 74.4 and the specific 
 heat of the metal is 0.031. Calculate the atomic weight of lead. 
 
CHAPTER XXX. 
 ELEMENTS OF GROUP V. 
 
 Family M. 
 
 (Vanadium, V. 51.2) 
 (Columbium,Cb. 94. ) 
 (Tantalum, Ta. 181. ) 
 
 The elements of this family are lustrous gray solids which are very rare 
 and difficult to isolate. Vanadium is the least uncommon. The mem- 
 bers of the family are closely related to one another. 
 
 Vanadium, which was first isolated by Roscoe in 1867, is found in the 
 complex mineral, vanadinite, Pb 4 (PbCl) (VO 4 ) 3 . It is slowly acted upon 
 by air at ordinary temperatures, but when heated it burns brilliantly, 
 forming the reddish-brown vanadium pentoxide, V 2 O 5 . This oxide in- 
 teracts with bases giving vanadaies. The metal possesses very feeble 
 base-forming properties. Vanadium also combines with nitrogen at red 
 heat to form the yellowish-red vanadium nitride, VN. The following 
 are among the more important of its compounds: V 2 O, V 2 O 2 , V 2 O 3 , 
 VO 2 , V 2 O 5 , VC1 2 , VC1 3 , VC1 4 , VOC1 3 , VOC1 5 , V 2 S, V 2 S 3 , and H 3 VO 4 . 
 The element is frequently prepared by heating VC1 2 in a stream of 
 hydrogen. 
 
 Columbium (niobium) and tantalum are found in the rare minerals 
 columbite, (Mn,Fe) (Cb,Ta) 2 O 6 , and tantalite, (Fe,Mn)Ta 2 O 6 , respec- 
 tively. These two elements likewise possess feebly base-forming properties. 
 The chief compounds are the columbates and tantalates. Other com- 
 pounds are, Cb 2 O 2 , Cb 2 O 4 , Cb 2 O 5 , CbOCl 3 , CbCl 3 , CbCl 5 , H 3 CbO 4 , 
 Ta 2 O 4 , Ta 2 O 5 , TaCl 5 and H 3 TaO 4 . 
 
 343 
 
CHAPTER XXXI. 
 ELEMENTS OF GROUP VI. 
 
 Family M. 
 
 Chromium, Cr. 52.1 
 (Molybdenum, Mo. 96.0)- 
 (Tungsten, W. 184.0) 
 (Uranium, U. 238.5) 
 
 The elements of this group possess many properties in common. They 
 are hard metals possessing a metallic luster and a high specific gravity. 
 They are attacked by acids, but are not acted upon by air and water at 
 ordinary temperatures. They are readily oxidized or redaced and their 
 corresponding compounds are structurally and chemically similar. Their 
 higher oxides are acid anhydrides, while their lower oxides are basic in 
 character. They form two classes of salts, the "ous" and the "ic"; 
 also two classes of compounds, "ites" and "ates," in which the elements 
 appear in the negative ions. 
 
 The maximum valence shown by the elements of this group is VI. 
 Chromium, however, in perchromic acid, H 2 Cr 2 O 8 , reaches a valence of 
 VII. 
 
 CHROMIUM, Cr. 
 At. Wt. 52.1 Sp. Gr. 6.92 
 
 Chromium does not occur in nature in the uncombined condition, and 
 its natural compounds are neither abundant nor widely distributed. The 
 chief source of chromium is chrome iron ore, or chromate, FeCr 2 O 4 . It 
 is prepared most conveniently by the " Goldschmidt" process. Chro- 
 mium prepared by this method is " passive " and does not displace hydrogen 
 from hydrochloric acid until it is warmed with the acid. When removed 
 from the acid and left in the air, it changes slowly into the inactive form 
 again. Chromium fuses in the electric arc, but not in the oxyhydrogen 
 flame. 
 
 Chromium forms the following oxides and hydroxides: hypothetical 
 chromous oxide, CrO, and chromous hydroxide, Cr(OH) 2 , which are 
 distinctly basic; chromic oxide, Cr 2 O 3 , and chromic hydroxide, Cr(OH) 3 , 
 which are weakly basic; chromic oxyhydroxide, HCrO 2 , which is acidic; 
 chromium trioxide, Cr 2 O 3 , a water solution of which is called chromic acid, 
 H 2 CrO 4 .* 
 
 It is obvious that chromium gives four classes of compounds, chromous 
 and chromic salts which correspond to Cr(OH) 2 and Cr(OH) 3 , and 
 * The compound H 2 CrO 4 has not been isolated. 
 344 
 
ELEMENTS OF GROUP VI. 345 
 
 chromites and chromates, to HCrO 2 and H 2 CrO 4 , respectively. There 
 is also another class, perchromates, which correspond to perchromic acid, 
 H 2 Cr 2 O 8 . 
 
 The chromous salts are not stable in the air owing to the ease with 
 which they pass into the condition of higher oxidation of chromic com- 
 pounds. All of the chromic salts are hydrolyzed, CrCl 3 and Cr 2 (SO 4 ) 3 to 
 a small degree, while neither the carbonate nor sulphide is stable in water. 
 (In this respect chromic salts are very similar to corresponding aluminum 
 salts.) The perchromates are unstable at ordinary temperatures. 
 
 It is to be noted that those chromium ions of " lower valence, or with 
 the smaller electrical charge, are basic; but as the valence increases, or 
 as the amount of electrical energy which they carry increases, the basic 
 property becomes less and less, and acidic properties begin to manifest 
 themselves."* 
 
 Experiments I. Properties of Chromium. 
 
 Examine the metal and note its most obvious physical properties. 
 Does it tarnish in the air? 
 
 Experiment II. Compounds of Chromium. 
 
 (a) Examine the compounds of chromium (end shelf). Give the color 
 and structural formula of each compound. Record the valence of the 
 chromium atom in each compound, and state whether it behaves as an 
 acid-forming or base-forming element. Tabulate the foregoing data. 
 
 (b) What is the formula, color and electrical charge of each of the 
 following ions: the dichromate ion and the chromic ion. Hint: 
 Examine solutions of the salts giving these ions (end shelf). 
 
 Experiment III. Preparation of a Chromate. 
 
 Mix 5 grams of potassium carbonate with equal amounts of potassium 
 nitrate (supplies oxygen) and potassium hydroxide in an iron crucible 
 (T.O.). Heat the mixture at low temperature until it melts, then stir 
 (use iron rod) in 5 grams of powdered chromite. Heat strongly in the 
 flame of a blast-lamp until further heating produces no more change. 
 Allow the crucible to cool, then dissolve the contents in a little boiling 
 water. Filter. What is the color of the solution ? The color is due to 
 the presence of what ion ? 
 
 4 (FeO,Cr 2 3 ) + 8(K 2 O,CQ 2 ) + yO 2 -> 8(K 2 O,CrO 3 ) + 2 Fe 2 O 3 -f 8CO 2 . 
 
 Experiment IV. Formation of a Dichromate from a Chromate. 
 
 To 5 cm. 3 of potassium chromate solution add sulphuric acid, drop by 
 drop, until the yellow color changes to orange. The color is due to the 
 presence of what ion? Evaporate in small beaker to crystallizing point; 
 cool slowly. Note color and shape of crystals. 
 
 * Jones, Elements of Inorganic Chemistry 
 
34^ EXPERIMENTAL CHEMISTRY. 
 
 Interpret the following equations: 
 
 K 2 CrO 4 + H 2 CrO 4 K 2 Cr 2 O 7 + H 2 O. 
 2K 2 CrO 4 + H 2 SO 4 K 2 Cr 2 O 7 + K 2 SO 4 + H 2 O. 
 
 Experiment V. Formation of a Chromate from a Bichromate. 
 
 To 5 cm. 3 of potassium dichromate solution add enough sodium hy- 
 droxide solution to turn the color yellow. The color is due to the presence 
 of what ion ? What salt has been formed by the interaction ? Equation ? 
 Evaporate until a crust forms, then set aside to crystallize. Record color 
 of crystals. 
 
 Experiment VI. Formation of Chromium Trioxide (Chromic Anhy- 
 dride); its Oxidizing Power. Chromic Acid.* 
 
 Prepare about 10 cm. 3 of a warm saturated aqueous solution of potas- 
 sium dichromate in a beaker. Filter the solution if it is not clear. Now 
 add carefully, a drop at a time, an equal volume of concentrated sul- 
 phuric acid. Red needle-shaped crystals of chromium trioxide will 
 separate.* Equation? When the mixture has cooled, filter through a 
 plug of glass wool or asbestos (not filter paper). (Caution: Do not 
 allow the crystals to come in contact with the hand.} By means of a glass 
 or porcelain spatula remove a few of the crystals and place them upon a 
 piece of filter paper. Account for the result. Treat a few of the crystals 
 in a test tube with a little hydrochloric acid and warm gently. 
 
 Identify the gas liberated (Hint: Odor?). Explain. What is the 
 green substance formed which remains in solution? 
 
 2CrO 3 + i2HCl 3 -> 2 CrCl 3 + 30, + 6H 2 O. 
 
 Pour a drop or two of alcohol upon some of the crystals. The action 
 (oxidation) is so violent that the alcohol frequently takes fire. Try the 
 solubility of some of the crystals in water ( ?). What is formed ? Equa- 
 tion? 
 
 The oxidizing power of chromates and dichromates in the presence of 
 an acid is due to the presence of what compound? 
 
 Experiment VII. Formation of Perchromic Acid. 
 
 To 5 cm. 3 of a solution of potassium chromate add sulphuric acid until 
 the yellow color of the solution changes to orange (avoid large excess of 
 acid), then add hydrogen peroxide. If the procedure has been successful, 
 blue color will be imparted to the solution by the unstable perchromic acid, 
 H 2 Cr 2 8 . 
 
 * If the solutions are dilute and there is not present a body capable of being oxid- 
 ized the decomposition is limited to the liberation of chromic acid, H 2 O.Cr O 3 . 
 
ELEMENTS OF GROUP VI. 347 
 
 Note. The perchromic acid is supposed to be a compound of CrO 3 
 and H 2 O 2 . 
 
 K 2 CrO 4 + H 2 SO 4 -> (H 2 O.CrO 3 ) + K 2 SO 4 . 
 
 ^ (H 2 O.CrO 3 ) -> K 2 Cr 2 O 7 .+ H 2 O. 
 2 K 2 Cr~6 4 + lp"O 4 - K 2 Cr 2 O 7 + K 2 SO 4 + H 2 O, 
 
 K 2 Cr 2 7 + H 2 S0 4 K 2 S0 4 + 2 CrO 3 + H 2 O 
 2 Cr0 3 + H 2 2 H 2 Cr 2 O s . 
 
 The above reaction is used as a delicate test for either hydrogen perox- 
 ide or a chromate. 
 
 Experminent VIII. Chromium as a Base-forming Element. Re- 
 duction of a Chromate to a Chromic Salt. Chrome Alum. 
 
 To a cold saturated solution of potassium dichromate acidified with 
 sulphuric acid, add any one of the following reducing agents until the 
 reddish-yellow color has changed to a dark violet: SO 2 (sulphurous 
 acid), H 2 C 2 O 4 . 2 H 2 O (oxalic acid), C 2 H 5 OH (alcohol). 
 
 The solution now contains chromium sulphate, Cr 2 (SO 4 ) 3 . The color 
 is due to the presence of what ion ? Set the solution aside and allow it to 
 evaporate spontaneously. Observe the color and form of the crystals. 
 Compare them with crystals of chrome alum, K 2 SO 4 ,Cr 2 (SO 4 ) 3 .24H 2 O. 
 Your conclusions as to the identity of the crystals ? The changes may be 
 represented thus: 
 
 2 Cr0 3 + 3 S0 2 -> Cr 2 (S0 4 ) 3 , 
 
 K 2 Cr 2 7 +H 2 O.S0 2 + H 2 SO 4 Cr 2 (SO 4 ) 3 + K 2 SO 4 + 4 H 2 O. 
 K 2 Cr 2 7 + 4 H 2 S0 4 -> Cr 2 (S0 4 ) 3 + K 2 S0 4 + ( 3 O) + 4 H 2 O, 
 3 C 2 H 5 OH + ( 3 0) -> 3 C 2 H 4 + 3 H 2 0. 
 
 (Aldehyde) 
 K 2 SO 4 + Cr 2 (SO 4 ) 3 + 2 4 H 2 O (K 2 SO 4 ,Cr 2 (SO 4 ) 3 .2 4 H 2 O) 
 
 Chrome alum. 
 
 Is chrome alum a double or complex salt ? Give reasons for your an- 
 swer. Does chromium function in this salt as an acid or base-forming 
 element ? Potassium chrome alum is the analogue of potassium aluminum 
 sulphate ("alum"). How can a chromate be changed to a dichro- 
 mate ? A dichromate to a chromate ? A chromate to a chromic salt ? 
 
 Experiment IX. Hydrolytic Decomposition of Chromium Sulphate. 
 
 Place 5 cm. 3 of a solution of chromium sulphate (chrome alum will do) 
 in a test tube and warm gently. Note the transition from the violet 
 chromic sulphate to the green chromic sulphate. This transition has 
 been attributed to hydrolytic decomposition: 
 
 2 Cr 2 (S0 4 ) 3 + H 2 0^ (Cr 4 0(S0 4 ) 4 )S0 4 + H 2 SO 4 . 
 Experiment X. Formation of Chromyl Chloride. A Test for a Chloride. 
 
 In a test tube fitted with a delivery tube warm a mixture of a chloride 
 and potassium dichromate with strong sulphuric acid. Pass the red- 
 
348 EXPERIMENTAL CHEMISTRY. 
 
 brown vapor (chromyl chloride, CrO 2 Cl 2 ) which is disengaged into a 
 second test tube containing a solution of an alkaline hydroxide. The 
 formation of a chromate is indicated by the yellow color which the solution 
 assumes, and may be confirmed by acidifying and adding a lead-nitrate 
 solution. Presence of the chromate is proof of the presence of a chloride 
 in the original mixture. As no corresponding bromine and iodine com- 
 pounds are known, it is obvious that by means of this test it is possible 
 to detect a chloride in the presence of either a bromide or iodide. 
 
 K 2 Cr 2 7 + 4 KC1 + 3 H 2 S0 4 2 Cr0 2 Cl 2 + 3 K 2 SO 4 + 3 H,O. 
 
 CrO 2 Cl 2 + 4NH 4 OH (NH 4 ) 2 CrO 4 + 2NH 4 C1 + 2H 2 O. 
 
 Experiment XI. Formation of Chromic Hydroxide. 
 
 To 10 cm. 3 of a solution of chrome alum add ammonium hydroxide 
 in slight excess. Result ? Equation ? Wash the precipitate of chromic 
 hydroxide, Cr(OH) 3 , with hot water; dry and reserve it for the following 
 experiment. 
 
 Experiment XII. Oxidation of a Chromic Compound to a Chromate. 
 
 (a) Mix a portion of the chromic hydroxide prepared in the foregoing 
 experiment with equal portions of potassium nitrate and sodium car- 
 bonate and fuse the mixture in an iron crucible or on a piece of platinum 
 foil. Dissolve the fused mass in water and filter. What is the color of 
 the filtrate ? Test it for potassium chromate by acidifying with acetic 
 acid and adding a solution of lead nitrate. Result? 
 
 (b) To 5 cm. 3 of a solution of a chromium salt (chrome alum) add 
 sufficient sodium hydroxide to make solution alkaline. How add a 
 large volume of bromine water and heat gently. What is the nature of 
 the change? Lead peroxide and other oxidizing agents may be sub- 
 stituted for the bromine water. 
 
 Experiment XIII. Analytical Reactions. 
 
 (a) Make a borax bead, touch it with a small quantity of any chromium 
 compound and heat in both the oxidizing and reducing flames. The 
 grass-green color is imparted to the bead by all of the chromium com- 
 pounds. 
 
 (b) Reactions of chromates. 
 
 1. Recall the action of sulphuric acid on a solution of potassium 
 chromate. Equation? 
 
 2. Boil a potassium dichromate solution which has been acidified 
 with dilute acid with i cm. 3 of alcohol. Account for results. Equation ? 
 
 3. Pass hydrogen sulphide into an acid solution of potassium chro- 
 mate. Results ? Does the hydrogen sulphide act as an oxidizing or re- 
 ducing agent? Equation? 
 
 4. Recall the action of hydrogen peroxide on an acid solution of a 
 chromate. Equation ? 
 
 5. Add a solution of barium chloride to a potassium chromate solu- 
 tion (?). Equation? 
 
ELEMENTS OF GROUP VI. 349 
 
 6. Repeat 5, using a potassium dichromate solution. Compare 
 results. Equation ? 
 
 7. Repeat 5 and 6, using a solution of lead nitrate or lead acetate. 
 Do dichromates precipitate dichromates or chromates when added to 
 a solution of a salt whose chromate is insoluble? 
 
 (c) Reactions of chromic salts. 
 
 1. Recall the effect of boiling a solution of a chromic salt (see Exp. 
 IX). Equation? 
 
 2. To a solution of potassium chromium sulphate (chrome alum) 
 add slowly a little of a solution of sodium hydroxide ( ?) ; then in excess ( ?) 
 Equations? Does aluminum hydroxide behave in a manner similar 
 to the chromium hydroxide ? Equation ? Is the chromium hydroxide 
 reprecipitated when the solution is boiled? 
 
 3. Treat a solution of chrome alum with sodium carbonate solution. 
 Account for the precipitation of chromium hydroxide. Ionic equations ? 
 
 4. Add ammonium sulphide to another portion of the solution of 
 chrome alum. Filter, and wash the precipitate until free from odor. 
 Prove that it is not a sulphide. What other metal fails to form a stable 
 sulphide in the presence of water ? Equation ? 
 
 (d) Ions. What ions are yielded by chromates and chromic salts? 
 
 MOLYBDENUM, MO. 
 
 At. Wt. 96.0 Sp. Gr. 8.6. 
 
 Molybdenum is found chiefly in the two uncommon minerals, molyb- 
 denite, MoS 2 , and wuljenite, PbMoO 4 . It is a hard, silver-w r hite metal 
 which is less fusible than platinum. At ordinary temperatures, it re- 
 mains unchanged in the air, but on heating it oxidizes to molybdenum 
 trioxide, MoO 3 . This oxide interacts with the alkalies forming molyb- 
 dates of the type of sodium molybdate, Na 2 MoO 4 .ioH 2 O. Nitric acid 
 precipitates a hydroxide MoO(OH) 4 from solutions of these molybdates. 
 This hydroxide gives molybdic acid, H 2 MoO 4 , when heated sufficiently 
 to dry it. When sodium phosphate is added to a solution of ammonium 
 molybdate in nitric acid the yellow pulverulent ammonium pliosphomolyb- 
 date, (NH 4 ) 3 PO 4 .iiMoO 3 .6H 2 O, is precipitated. This is 'one of the 
 more familiar of the complex phosphomolybdates. The formulae of some 
 of the more important compounds of molybdenum are as follows: MoO, 
 Mo 2 O 8 , MoO 2 , MoO 3 , MoCl 2 , MoCl 3 , MoCl 4 , MoCl 5 , MoO 2 Cl 2 , MoOCl 4 , 
 MoS 2 , MoS 3 , Na 2 MoO 4 , K 2 Mo 3 O 13 .ioH 2 O, Na 2 Mo 4 O 13 .6H 2 O. 
 
 It is evident that molybdenum possesses valences ranging ordinarily 
 from II to VI. 
 
 TUNGSTEN, W. 
 
 At. Wt. 184.0 Sp. Gr. 19.1. 
 
 Tungsten occurs in the minerals wolframite, 3FeWO 4 .3MnWO 4 , 
 scheelite, CaWO 4 , and hubnerite, MnWO 4 . The metal possesses prop- 
 
350 EXPERIMENTAL CHEMISTRY. 
 
 erties similar to those of molybdenum. If forms a class of compounds, 
 many of which are analogues of the compounds of the previously men- 
 tioned element. The salts corresponding to tungstic acid, H 2 WO 4 , are 
 known as tungstates. The following list contains the formulae of some 
 of the better known compounds: WO 2 , WO 3 , WC1 2 , WC1 4 , WC1 5 , WCl e , 
 
 The metal is used in the manufacture of tungsten steel, a very hard 
 variety of which contains about 5 per cent, of tungsten. 
 
 URANIUM, U. 
 
 At. Wt. 238.5 Sp. Gr. 18.7. 
 
 Uranium is found in the ore pitchblende, which contains the mineral 
 uraninite, U 3 O 8 , and several other rare minerals. It is a heavy, silvery- 
 white metal which decomposes water at ordinary temperatures and burns 
 in the air at 175 when in the powdered form. Some of the more im- 
 portant compounds are, uranous oxide, UO 2 , uranic anhydride, UO 3 , 
 uranous chloride, UC1 4 , uranic or uranyl sulphate, UO 2 SO 4 .6H 2 O, 
 sodium diuranate, Na^C^Oy. Many of the uranyl compounds are yel- 
 low in color, with green fluorescence. Sodium diuranate is used in mak- 
 ing uranium glass which shows a yellowish-green fluorescence. 
 
 The compounds of uranium are of particular interest because of being 
 the object of much investigation in connection with the phenomena of 
 radio-activity. In 1898, Becquerel noticed that all compounds of uran- 
 ium gave out radiation capable of affecting a photographic plate covered 
 with dark light-proof paper. His observations eventually led to the 
 discovery, by Mme. Curie of the element radium. (The student is urged 
 to consult some text-book on the subject of radio-activity, for example, 
 Rutherford, Radio-activity.) 
 
 PROBLEMS. 
 
 1. Calculate the percentage composition of (a) lead chromate, 
 PbCrO 4 , (b) chromite, FeCr 2 O 4 . 
 
 2. If 10 grams of chromous chloride, CrCl 2 , yield 5.75 grams of chlor- 
 ine, what is the atomic weight of chromium? 
 
 3. 10 grams of silver chloride are formed by the interaction of silver ni- 
 trate and 3.6865 grams of chromic chloride; what is the atomic weight 
 of chromium? 
 
 3 AgN0 3 + CrCl 3 Cr(N0 3 ) 3 
 
CHAPTER XXXII. 
 
 ELEMENTS OF GROUP VII. 
 
 Family M. 
 MANGANESE, Mn. 
 
 In the periodic table, according to the present classification, man- 
 ganese stands alone on the left side of the eighth column. The elements 
 of the halogen group occupy the right side. It is both an acid-forming 
 and base-forming element, and probably forms as large a variety of 
 compounds as any element known. This latter property is due to the 
 many degrees of valence which manganese can manifest. 
 
 Although manganese occurs native in small amounts associated with 
 iron in meteorites, it is found principally in combination in minerals, 
 some of which are rather widely distributed. The chief source of the 
 metal is pyrolusite, MnO 2 . Other minerals containing it are: baunite, 
 Mn 2 O 3 , hausmannite, Mn 3 O 4 , the hydrated form, manganite, MnO(OH), 
 manganese spar, MnCO 3 , and manganese blende, MnS. 
 
 Manganese may be obtained by heating the oxides with carbon in 
 an electric furnace or by electrolysis of the fused chloride. It is prepared 
 more conveniently, however, by the " Goldschmidt Process"; i. e., by 
 mixing the oxide with finely divided aluminum and igniting the mixture. 
 The aluminum takes the oxygen and sets free the manganese. 
 
 Manganese is a hard, grayish-white, brittle metal of brilliant luster. 
 It is slightly magnetic and fuses at 1900. It is permanent in dry air, 
 but readily oxidizes superficially on exposure to moist air, and slowly 
 decomposes boiling water with the evolution of hydrogen when the 
 metal is in the finely divided condition. It dissolves readily in dilute 
 acids with the formation of manganous salts. The metal is used as a com- 
 ponent of the alloys, ferro-manganese (20 to 75 per cent, manganese), 
 and spiegeleisen (iron-manganese carbide), which are of use in the metal- 
 lurgy of Bessemer steel. " Manganin" is an alloy of copper, manganese 
 and nickel, containing from eight to ten per cent, of manganese and three 
 to four per cent, of nickel. In the form of wire it is much used in the 
 construction of resistance coils because of its very low resistance tem- 
 perature coefficient. 
 
 Manganese forms five rather well defined sets of compounds which 
 correspond to its series of oxides. The composition of these compounds 
 is as follows: 
 
 Manganous, 
 MnO 
 Mn(OH) 2 
 MnSO 4 
 MnCl 2 
 etc. 
 
 Manganic, 
 
 Mn 2 O 3 
 
 Mn(OH) 3 
 
 Mn 2 (S0 4 ) 3 
 
 (MnCl a ) 
 
 Manganites, 
 MnO 2 
 H 2 MnO 3 
 CaMnO 3 
 
 etc. 
 
 Manganates, 
 MnO 3 
 H 2 MnO 4 
 K 2 MnO 4 
 etc. 
 
 Permanganates. 
 Mn 2 O 7 
 HMnO 4 
 KMnO 4 
 etc. 
 
352 EXPERIMENTAL CHEMISTRY. 
 
 The manganous salts are pale pink in color and are but slightly hydro- 
 lyzed. The oxide is green powder and a strong base. The manganic 
 salts are violet in color and are completely hydrolyzed. Manganese 
 sesquioxide is a weak base. The manganites are usually of dark color 
 and are strongly hydrolyzed. Manganese dioxide is a black solid and 
 behaves as an indifferent oxide. The manganates are green in color 
 and are very easily hydrolyzed, the free acid decomposing and yielding a 
 higher acid (HMnO 4 ) and a lower oxide (MnO 2 ). 
 
 Manganese trioxide is an amorphous red solid soluble in w r ater. It is 
 viewed as the anhydride of manganic acid, H 2 MnO 4 . The permanganates 
 are purplish-red in color and are not hydrolyzed by water. The sept- 
 oxide of manganese, Mn 2 O 7 , is an oily, dark liquid which is presumably 
 the anhydride of permanganic acid, HMnO 4 . 
 
 Manganese has a valence of II, III, IV, VI and VII. If the existence 
 of manganese tetroxide, MnO 4 , is admitted, then the metal has a maxi- 
 mum valence of VIII. 
 
 It should be noted that as the valence of manganese increases, its 
 basic nature diminishes, and that it loses all of its basic nature and 
 manifests strong acid properties w r hen in the condition of higher valence. 
 
 MANGANESE, MR. 
 
 At. Wt. 55.0 Sp. Gr. 7.2-8.0. 
 
 Experiment I. Note the most obvious physical properties of man- 
 ganese. 
 
 Does it oxidize in the air? What are some of the uses of manganese? 
 Experiment II. Compounds of Manganese. 
 
 Follow directions given in Experiment II, " Chromium," substituting 
 the word permanganate for the word dichromate. 
 
 Experiment III. Preparation of a Manganate. Oxidation of a 
 Manganate to a Permanganate. 
 
 (a) Grind in a mortar 5 grams of potassium hydroxide and 2.5 grams 
 of potassium chlorate; transfer the mixture to an iron crucible (T. O.) 
 or a porcelain crucible, and heat until the mixture fuses, then add gradu- 
 ally 5 grams of powdered manganese dioxide while stirring with an iron 
 rod or the reverse end of a file. Maintain the crucible at a red heat for 
 15 to 20 minutes. Dissolve the green mass (potassium manganate) in a 
 little cold water, then decant the clear liquid away from the residue. 
 What is the color of the solution? This color is due to the presence 
 of what ion? 
 
 2 MnO 2 + 6KOH + KC1O 3 -> 3 K 2 MnO 4 + KC1 + 3H 2 O- 
 
 (b) Divide the green solution from (a) into four parts. Dilute one 
 portion with an equal volume of distilled water, and set aside for several 
 
ELEMENTS OF GROUP VII. 353 
 
 days. It will become purple in color owing to the oxidation of the 
 manganate to the permanganate. Pass carbon dioxide (generator) 
 through one portion of the manganate solution. Result? 
 
 3 K 2 MnO 4 + 2CO 2 -> 2 KMnO 4 + MnO 2 + 2K 2 CO 3 . 
 
 What is the color of the permanganate ion? Interpret the following 
 equations: 
 
 K 2 MnO 4 + 2H 2 O -> 2KOH + (H 2 MnO 4 ), (i) 
 
 or 2 K' + MnO" 4 + 2H' + 2 OH'-> 2 K' + 2 OH' + (H 2 MnO 4 ), 
 
 H 2 MnO 4 = H 2 O,MnO 3 , 
 
 3(H 2 O,MnO 3 ) H 2 O,Mn 2 O 7 + MnO 2 + 2H 2 O, (2) 
 or H 2 O,Mn 2 O 7 = 2 HMnO 4 
 
 3 H 2 MnO 4 2 HMnO 4 + MnO 2 + 2H 2 O (3) 
 2 KOH + 2 HMnO 4 2 KMnO 4 + 2H 2 O 
 
 The partial equations may be summarized: 
 
 3 K 2 MnO 4 + 2H 2 O 2 KMnO 4 + 4 KOH + MnO 2 
 
 The ionic equation: 
 
 6K' + 3MnO" 4 + 2H' + 2 OH'-> 6K' + 2MnO / 4 + Mn O 2 + 4OH'. 
 
 What is the difference between the manganate-ion and the perman- 
 ganate-ion ? 
 
 Experiment IV. Oxidizing Power of Manganates and Permanganates. 
 
 Note. Potassium permanganate as well as potassium manganate 
 are powerful oxidizing agents. 
 
 (a) Manganates. To a warm dilute solution of oxalic acid add 
 an alkaline solution of potassium manganate drop by drop. The oxalic 
 acid is oxidized to carbon dioxide (prove that it is evolved) and water. 
 The manganese is reduced to the bivalent condition with the formation 
 of manganous sulphate, MnSO 4 . 
 
 (b) Permanganates. 
 
 1. Repeat (a) with a solution of potassium permanganate to which 
 has been added an equal volume of dilute sulphuric acid. Result? 
 Equation ? 
 
 2. To 5 cm. 3 of sulphurous acid add an acidified solution of potassium 
 permanganate drop by drop. Results ? Equation ? 
 
 3. Add sulphuric acid to a solution of ferrous sulphate, FeSO 4 , 
 then add potassium permanganate solution drop by drop. The ferrous 
 sulphate is oxidized to ferric sulphate, Fe 2 (SO 4 ) 3 , and the permanganate 
 is reduced. Explain the color changes. Equation? 
 
 4. Add a solution of potassium hydroxide to a potassium perman- 
 ganate solution. Result? Equation? 
 
 5. Recall or again try the action of hydrogen peroxide" on an acidified 
 solution of potassium permanganate. 
 
354 EXPERIMENTAL CHEMISTRY. 
 
 Interpret the following equations: 
 
 2 KMnO 4 + 3H 2 SO 4 -*K 2 SO 4 + 2MnSO 4 + 3H 2 O + (50). 
 
 5 C 2 H 2 4 + (50) -r- ioC0 2 + 5 H 2 0. 
 2 FeS0 4 + H 2 S0 4 + (O) Fe 2 (SO 4 ) 3 + H 2 O. 
 
 H 2 SO 3 + (O) H 2 SO 4 
 
 4 KMnO 4 + 4KOH 4 K 2 MnO 4 +_2_H a O + O 3 
 2KMnO 4 + (heat) K 2 MnO 4 + O 2 . 
 
 Experiment V. Reactions of Manganous Salts. 
 
 Note. Use a solution of manganous chloride, MnCl 2 , in performing 
 the following experiments. 
 
 (a) What color is imparted to the borax bead by manganese com- 
 pounds when heated in the oxidizing flame ? In the reducing flame ? 
 
 (b) Mix a little manganous chloride with sodium carbonate and 
 potassium nitrate, and fuse on a platinum foil. What is the color of 
 the fused mass? Identify the substance (?). 
 
 (c) (i) To a dilute solution of manganous chloride add ammonium 
 hydroxide. Result? Equation? 
 
 (2) Repeat (i) using a solution of manganous chloride to which has 
 been added ammonium chloride. Result ? 
 
 (d) Repeat (c) (i), using sodium hydroxide solution in excess. Equa- 
 tion? Treat the manganous hydroxide with bromine water. 
 
 (e) Repeat (c), using a solution of sodium carbonate or ammonium car- 
 bonate. Result? Prove that the precipitate is a carbonate (?). Equa- 
 tions ? 
 
 (/) Add ammonium sulphide to a solution of manganous chloride. 
 Prove that the precipitate is a sulphide (?). Equations? 
 
 (g) Pass hydrogen sulphide (generator in hood) through a solution 
 of manganous chloride to which has been added i cm. 3 of acetic acid. 
 Account for the negative results. 
 
 (h) What is the color of the manganous ion ? Of the manganate ion ? 
 Of the permanganate ion? 
 
 PROBLEMS. 
 
 i. A liter of a solution of potassium permanganate contains 10 grams 
 of the salt. How many grams of ferrous sulphate, FeSO 4 , can this 
 solution oxidize to ferric sulphate, Fe 2 (SO 4 ) 3 , in the presence of the 
 proper amount of sulphuric acid? 
 
CHAPTER XXXIII. 
 TRANSITION ELEMENTS. 
 
 FIRST LONG PERIOD (IRON ELEMENTS). 
 
 Iron, Fe. 55.9 
 Cobalt, Co. 59.0 
 Nickel, Ni. 58.7 
 
 The iron elements are among the most important elements tech- 
 nically, and are most interesting from the chemical stand-point. 
 
 These three elements, iron, cobalt and nickel, stand in a different 
 relation to one another than the members of the other eight groups; i.e., 
 they are not the corresponding members of successive periods as are 
 the elements of the families. They belong to the same period and form a 
 transition group between the first and second series of the first long 
 period. (See Periodic Classification.) These three elements are 
 closely related; in nature they are usually associated; the free metals 
 are magnetic and possess many physical properties in common. They 
 exhibit, however, a gradual transition in their chemical properties. 
 Thus, iron forms ferrates, M' 2 FeO 4 , and two basic oxides, ferrous oxide, 
 FeO, and ferric oxide, Fe 2 O 3 , each of which yields a series of stable salts 
 of the type, FeCl 2 (ferrous chloride) and FeCl 3 (ferric chloride), respec- 
 tively. Cobalt forms cobaltous and cobaltic salts, like CoCl 2 and Co 2 (SO 4 ) 3 . 
 Many of the cobalt^ salts (except the double salts) are unstable. Nickel 
 forms only one series of salts, namely, mckelous salts, like NiCl 2 . 
 
 It is obvious that these metals are related, on the one hand, through iron, 
 to chromium and manganese (recall such compounds as chromates, 
 manganate and ferrates), and on the other hand, through nickel, to copper 
 and zinc, both of which are bivalent elements and follow in the period. 
 
 IRON, Fe. 
 At. Wt. 55.9 Sp. Gr. 7.78. 
 
 Iron is one of the most abundant and widely distributed elements, 
 yet it is found native in only small amounts in meteorites and certain 
 volcanic ejects. Most of the rocks contain compounds of iron, and 
 many of the red and yellow soils owe their color to the presence of iron 
 compounds. Mineral waters holding iron compounds in solution are 
 known as "chalybeate waters." Minute .quantities of compounds of 
 iron are also found in chlorophyl and in the haemoglobin of the blood.* 
 
 * Ammonium sulphide interacting with the iron compounds present in the tissues 
 blackens the skin 
 
 355 
 
356 EXPERIMENTAL CHEMISTRY. 
 
 The chief ores of iron are red hematite and specular iron ore, Fe 2 O 3 (ores 
 found in the Lake Superior region, particularly in northern Michigan, 
 in Alabama, in Missouri and other regions of the United States) ; limonite 
 or brown hematite, 2Fe 2 O 3 ,3H 2 O (ores found chiefly in Alabama and 
 several other Southern States); magnetite or magnetic iron ore (loadstone), 
 Fe 3 O 4 (ores found in Pennsylvania, New York, Michigan and New 
 Jersey); jranklinite, similar to magnetite; and siderite or spathic iron ore, 
 FeCO 3 . Iron is also found in combination w T ith sulphur as iron pyrites 
 (Fool's gold),FeS. This pyrite is used in the manufacture of sulphuric 
 acid. Chalcopyrite, a sulphide of iron and copper, contains small quan- 
 tities of cobalt and nickel. 
 
 Iron ores are usually reduced by heating them with carbon and a flux 
 in a blast- furnace. The nature of the flux depends on the composition 
 of the iron ores; if the ores contain silica and clay, a flux of basic nature, 
 like limestone, is used; and, conversely, ores mixed with lime or magnesia 
 are heated with an acid flux, such as sand or clay-slate, in order that a 
 fusible slag may be formed. An impure iron (pig iron) is obtained by 
 this process. 
 
 (Discussions of the composition, properties and uses of cast iron, 
 wrought iron, Spiegel iron and steel will be found in the reference texts. 
 The student is urged to become familiar with the principles involved in 
 the manufacture of steel by such methods as the "Bessemer process," 
 the " Thomas-Gilchrist process," and the "Siemens-Martin process" 
 or " Open-hearth process.") 
 
 Pure iron may be prepared by reducing the oxide or oxalate in a 
 stream of hydrogen; electrolytic iron may be deposited from solutions 
 of certain salts by a properly regulated electric current; or the pure metal 
 may be prepared by Goldschmidt's process. 
 
 It is a white, lustrous metal, ductile and more malleable than wrought 
 iron. When finely divided it has a gray color. It is very tenacious; 
 and it is to this property, together with its abundance and ease with which 
 it can be prepared, that makes it the most valuable, industrially, of all 
 metals. Pure iron melts at 1800; wrought iron, at 1600; and cast iron, 
 noo-i3oo. At red heat it becomes soft and can be welded. It is 
 attracted to a magnet, but does not retain its magnetism. Iron is not 
 acted upon by dry air at ordinary temperatures, but in moist air it be- 
 comes coated with a "rust" which is probably a mixture of the oxide 
 and the hydroxide of iron, 2Fe 2 O 3 , (FeOH) 3 . It is not definitely under- 
 stood just how the product* is formed. The alkali hydroxides or car- 
 bonates prevent rusting. At red heat, massive iron decomposes water, 
 while finely divided iron decomposes water at 100. Iron dissolves in 
 dilute hydrochloric and sulphuric acids, hydrogen being evolved. Dilute 
 nitric acid dissolves iron with the formation of ferrous nitrate and am- 
 monium nitrate; concentrated nitric acid yields ferric nitrate and oxides 
 of nitrogen. When steel and cast iron, which contain iron, iron carbide, 
 Fe 3 C, and graphite are treated with cold dilute acids almost pure hydrogen 
 *The Corrosion of Iron, A. S. Cushman. Bui. No. 30, Dep't Agr. 
 
TRANSITION ELEMENTS. 357 
 
 is liberated, as the carbide and the graphite are not attacked; more con- 
 centrated acids, however, decompose the carbide, with the assistance 
 of nascent hydrogen, the freed carbon combines with hydrogen forming 
 hydrocarbons, which mix with the escaping gas, and give to it a very 
 disagreeable odor. With nitric acid, the carbide carbon behaves in a 
 slightly different manner. In this case the carbide passes into solution 
 in combination with hydrogen and oxygen. This compound imparts a 
 distinct color to the solution. Under similar conditions the depth of the 
 color varies with the proportion of carbon present. The graphite, how- 
 ever, is set free and separates out as with other acids. When a piece of 
 iron is immersed in very strong nitric acid (sp. gr. 1.42) for an instant 
 and then removed, it is not attacked by ordinary nitric acid and is in- 
 capable of displacing hydrogen and other elements lying below it in the 
 electromotive series. Iron exhibiting these peculiar properties is said 
 to be in the passive state. Such iron loses these properties and becomes 
 active when it is struck a sharp blow or is brought into contact with 
 ordinary iron. It was formerly supposed that the apparent chemical 
 inertness of the metal was due to the formation of a protecting layer 
 of oxide over its surface, but it has been shown recently to the satisfaction 
 of many that the passivity is due to an electrical condition of the metal.* 
 Nearly all of the compounds of iron are "ous" compounds in which 
 the metal is a bivalent positive ion, or "ic" compounds in which iron 
 is a trivalent positive ion. The oxides and hydroxides, FeO and Fe(OH) 2 , 
 Fe 2 O 3 and Fe(OH) 3 , are basic, the ferrous hydroxide being the stronger 
 base. The ferric salts derived from Fe(OH) 3 are hydrolyzed to con- 
 siderable extent. The ferrous salts are easily oxidized by the air to the 
 ferric condition. Only a few ferrates, K 2 FeO 4 ,BaFeO 4 , are known. 
 Such salts as potassium ferrocyanide, K 4 .Fe(CN) 6 and potassium ferri- 
 cyanide, K 3 .Fe(CN) 6 , yield complex anions which contain this element. 
 
 IRON, FC. 
 
 At. Wt. 55.9 Sp. Gr. 7.8. 
 Experiment I. Preparation of Iron by Reduction of Iron Oxide. 
 
 (a) Reduction by hydrogen. Assemble apparatus (Fig. 48) as described 
 in Chapter IX, Experiment VIII. Observe the precautions suggested. 
 Substitute iron oxide, Fe 2 O 3 , for the copper oxide in the porcelain boat. 
 Are the particles of iron oxide attracted by the magnet ? After the glass 
 tube has been maintained at a red heat for 10 to 15 minutes, remove 
 the heat and allow the boat and contents to cool in a stream of hydrogen. 
 What was formed in the anterior portion of the tube? Examine the 
 substance in the boat. What is its color? Are particles of it attracted 
 by the magnet? Expose it to the air (?). Why did you cool the sub- 
 
 * The Polarization Capacity of Iron and its Bearing on Passivity. Jour. Amer. 
 Chem. Soc., Nov., 1906. 
 
EXPERIMENTAL CHEMISTRY. 
 
 stance in an atmosphere of hydrogen? Identify the substance (?). 
 Equation ? What is " Venetian red " ? 
 
 (b) Goldschmidt's method (reduction with aluminum). Mix equal 
 volumes of powdered iron oxide (hammer scale) and aluminum powder, 
 and ignite (Caution!) the mixture on an iron plate or a brick. This 
 
 FIG. 48. (Smith and Keller.) 
 
 latter operation is performed most readily by inserting a piece of mag- 
 nesium ribbon in the mixture and setting fire to the projecting part. 
 Result ? Equation ? 
 
 Note. "Thermit" may be substituted for the above mixture. 
 
 Experiment II. Properties of Iron. 
 
 (a) Examine cast iron, wrought iron and steel. Note their most 
 obvious physical properties. Give the approximate composition of each. 
 Examine a piece of piano wire; it is about 99.7 per cent. pure. 
 
 (b) Recall the behavior of a piece of iron heated to redness when 
 plunged into a jar of oxygen. Equation? Recall or try the action of 
 iron when treated with dilute acids. Equations ? 
 
 Does iron " rust " in dry air ? In moist air ? Equation ? 
 
 Experiment III. Compounds of Iron. 
 
 (a) Examine the compounds of iron (end shelf). Give the color and 
 structural formula of each compound. Record the valence of the iron 
 atom in each compound, and state whether it manifests the properties 
 of an acid-forming or base-forming element. Tabulate the above data. 
 
 (b) What is the formula, color and electrical charge of the ferrous 
 ion, the ferric ion and the f erro-cyanide ion ? 
 
 Hint. Examine solutions of the salts which yield these ions (end 
 shelf). 
 
 Experiment IV. Preparation of Ferrous Sulphate. 
 
 Place about 20 grams of iron (free from rust) in the form of filings, 
 nails or wire, in an Erlenmeyer flask in which there should be inserted a 
 cork fitted with a glass jet to allow gas to escape. Pour on the iron 175 
 cm. 3 of dilute sulphuric acid, add a few drops of concentrated acid and 
 
TRANSITION ELEMENTS. 359 
 
 warm gently if the action appears slow. Note the odor of the escaping 
 gases ( ?). Account for the odor. Allow the action to continue for 10 or 
 15 minutes, adding sufficient strong acid to keep up a brisk action, while 
 other experiments are proceeded with. When nearly (but not quite) 
 all of the iron has dissolved, filter into a casserole containing 2 cm. 3 
 of sulphuric acid. Note the color of the solution and set it aside for a day. 
 Pour off the mother liquid from the crystals, and wash the latter with cold 
 water by decantation. Dry the crystals between sheets of filter paper. 
 Note their color, taste, and solubility in water (?). Put a few of them 
 into hard test tube and heat, gently at first, then strongly (?). Compare 
 with the corresponding properties of ferrous sulphate, FeSO 4 .7H 2 O, 
 crystals (end shelf). Identify the prepared crystals. Equation? Write 
 the structural formula for ferrous sulphate. What is the valence of iron 
 in this compound? 
 
 What is "green vitriol"? "oil of vitriol"? "copperas"? "white 
 viriol " ? iron protosulphate ? writing ink (black) ? 
 
 Experiment V. Preparation of Ferrous Ammonium Sulphate (Mohr's 
 Salt). 
 
 Weigh 5 grams of ferrous sulphate into an evaporating dish or a cas- 
 serole ; calculate the weight of and weigh out an equi-molecular quantity 
 of ammonium sulphate. Dissolve the salts separately in the smallest 
 volume of hot water. Add a few drops of sulphuric acid to the solution 
 of ferrous sulphate. Mix the solutions. Allow the mixture to cool 
 slowly and evaporate spontaneously (?). Describe the crystals and 
 compare them with those in the laboratory (end shelf). Collect the 
 crystals in a funnel, the stem of which is closed with a loose plug of 
 glass-wool (or better, provided with a platinum filter cone), allow the 
 mother liquid to drain off, then wash the crystals with a small quantity 
 of cold water and dry with filter paper. 
 
 Is "Mohr's Salt," (NH 4 ) 2 SO 4 ,FeSO 4 ,6H 2 O, a complex or compound 
 salt ? Reason for your answer ? Equation ? 
 
 Experiment VI. Preparation of Ferric Ammonium Sulphate, Iron- 
 Ammonium Alum. 
 
 Directions are the same as those given in Experiment VI, except that 
 ferric sulphate, Fe 2 (SO 4 ) 3 , is used instead of ferrous sulphate. Equation ? 
 
 To which class of salts, complex or compound, does iron-ammonium 
 alum, (NH 4 ) 2 SO 4 ,Fe 2 (SO 4 ) 3 ,2 4 H 2 O, belong? 
 
 Experiment VII. Hydrolysis of Ferric Salts. 
 
 Dissolve equal weights (about 0.5 gram) of ferrous sulphate (use 
 ferrous ammonium sulphate) and ferric sulphate, Fe 2 (SO 4 ) 3 , in equal 
 volumes of water in separate test tubes, then warm slightly. Observe 
 the color of each by looking down through the solution at a piece of white 
 paper (?). Test each solution with litmus paper (?). Add 2 or 3 cm. 3 
 of pure concentrated sulphuric acid to each solution and observe the 
 
360 EXPERIMENTAL CHEMISTRY. 
 
 colors again (?). What is the color of the ferrous ion? The ferric 
 ion is almost colorless?* Ferric hydroxide, Fe(OH) 3 , is a reddishb-rown 
 substance. Can you account for the change in color of the solution of 
 the ferric salt? Ionic equations? Which manifests the stronger basic 
 properties iron in the "ous" or iron in the "ic" condition? 
 
 Experiment VIII. Reactions of Ferrous and Ferric Salts. 
 
 (a) Make a borax bead and dissolve in it a small quantity of any 
 iron compound (preferably, the oxide); treat the bead successively with 
 the oxidizing (?) and reducing flames (?). Use a recently prepared 
 solution of ferrous sulphate (ferrous ammonium sulphate) and a dilute 
 solution of ferric chloride, FeCl 3 , for the following reactions. Treat a 
 portion of each, separately, with the following reagents. Compare the 
 two results obtained with each reagent. 
 
 (6) Ammonium hydroxide (?). Equations? Allow the tubes and 
 their contents to stand exposed to the air for 15 or 20 minutes. Result? 
 Equations ? 
 
 (c) Sodium hydroxide solution (?). Equation? Do the precipitates 
 dissolve in excess of the precipitant? 
 
 (d) Sodium carbonate ( ?). Prove by proper tests that the precipitates 
 do or do not contain the carbonic acid radical. Equations? 
 
 (e) A few drops of a solution of potassium ferricyanide, K 3 Fe(CN) 5 ( ?). 
 Equation ? What is " Turnbull's blue ? " 
 
 (/) A few drops of a potassium ferrocyanide, K 4 Fe(CN) 3 , solution (?). 
 Equation ? What is " Berlin or Prussian blue ? " 
 
 (g) Several drops of a potassium thiocyanate, KCNS, solution? 
 Equation ? Inasmuch as the ferric ion and the thiocyanate ion are color- 
 less, what is probably the source of the color? 
 
 (h) What reagents would you use to test for the presence of the ferrous 
 ion? For the ferric ion? 
 
 Note. Potassium ferrocyanide, potassium ferricyanide and potas- 
 sium thiocyanate belong to a class of substances known as " indicators"^ 
 By their use we are enabled to detect, for example, the presence of the 
 iron ion, and to determine also its state of oxidation; i.e., whether it is 
 in the "ous" or "ic" condition. The use of potassium thiocyanate 
 constitutes a very delicate test for detecting traces of the ferric ion. 
 
 Experiment IX. Ferrocyanides and Ferricyanides. 
 
 Using diluted solutions of potassium ferrocyanide, apply tests as given 
 in Experiment VIII, (ft), (c}, (d), to ascertain whether either the ferrous 
 
 * The ferric ion is almost colorless. The yellowish-brown color of solutions of 
 ferric salts is due to the piesence of ferric hydroxide (reddish-brown) produced by 
 hydrolysis. 
 
 t See Ostwald's Scientific Foundations of Analytical Chemistry. 
 
TRANSITION ELEMENTS. 361 
 
 or ferric ions are present or not. Result ? Is potassium ferrocyanide a 
 "complex" or a "compound" salt? Why? Indicate by an equation 
 how it probably dissociates. (See note, Chapter XXV. Experiment 
 XVII.) 
 
 Experiment X. Reactions of Iron Salts Involving Oxidations and 
 Reductions. 
 
 (a) Action of hydrogen sulphide on iron salts. 
 
 1. Saturate a dilute solution of a ferrous salt (ferrous-ammonium 
 sulphate) with hydrogen sulphide (?). Add ammonium hydroxide (?). 
 Filter, wash thoroughly the precipitate with cold water, and ascertain 
 whether it is a sulphide or hydroxide. Explain in terms of the ion- 
 product constant why hydrogen sulphide does not not precipitate all of 
 the iron from solutions of its ferrous salts. Ferric salts are reduced by 
 hydrogen sulphide. 
 
 2. Pass hydrogen sulphide into a dilute solution of ferric chloride (?). 
 Filter. Test the precipitate by burning it on a piece of filter paper (note 
 the odor) (?). Test the filtrate for the presence of ferrous ions (?). 
 
 2FeCl 3 + H 2 S 2FeCl 2 + 2HC1 + S., 
 2 Fe'" + S"-2Fe" +_S. 
 
 (b) Action of ammonium sulphide. 
 
 1. Ammonium sulphide precipitates black ferrous sulphide, FeS, 
 from solutions of ferrous salts. 
 
 2. Ferric salts are reduced by ammonium sulphide. The latter 
 produces a black precipitate of ferrous sulphide, soluble in hydrochloric 
 acid. 
 
 2 FeCl 3 + 3(NH 4 ) 2 S - 2 FeS + 6NH 4 C1 + S. 
 
 (c) Reduction of ferric salts to ferrous salts. 
 
 1. By nascent hydrogen. Place a few pieces of granulated zinc in a 
 small flask which is fitted with a cork provided with a glass jet for the 
 escape of gases. Pour 20 cm. 3 of ferric chloride into the flask and add an 
 equal volume of water. Now add sufficient hydrochloric acid to dissolve 
 the zinc with a rapid evolution of hydrogen. The action is hastened by 
 warming gently. Test a portion of the solution for the presence of ferrous 
 ions (?). Is the ferric ion present (test) ? Equations? 
 
 2. By stannous chloride. To 3 cm. 3 of a solution of ferric chloride 
 in a test tube add i cm. 3 of stannous chloride. Test portions of the 
 solution as in (a), i. Equations? 
 
 2FeCl 3 + SnCl 2 2FeCl 2 + SnCl 4 
 2 Fe"' + Sn" -* 2 Fe" + Sn"" 
 
 (d) Oxidation of ferrous salts to ferric salts. 
 
 i. By nitric acid. Into 5 cm. 3 of a hot dilute solution of ferrous 
 
362 EXPERIMENTAL CHEMISTRY. 
 
 sulphate, to which a little sulphuric acid has been added, pour 5 cm. 3 
 of concentrated nitric acid drop by drop. Test separate portions of the 
 liquid for the presence of ferrous and ferric ions (?). Make a record 
 of all tests and results. Equations? 
 
 2 HN0 3 ( 3 0) + 2 NO + H 2 0. 
 
 6FeS0 4 + 3H 2 S0 4 + 2 HNO 3 3 Fe 2 (SO 4 ) 3 + 4 H 2 O + 2 NO. 
 6FeCl 2 + 3HC1 + 2 HNO 3 6FeCl 3 + 4H 2 O + 2 NO. 
 
 2. By potassium bichromate in the presence of an acid. Pour 5 
 cm. 3 of dilute sulphuric acid into 15 cm. 3 of ferrous sulphate solution in 
 a small beaker. Now add by means of a burette a solution of potassium 
 bichromate drop by drop until a drop of the solution of iron salt trans- 
 ferred to white porcelain and tested with potassium ferricyanide solution 
 fails to give a blue color. This test shows the absence of what ion? 
 Test drops of the solution with a solution of potassium ferrocyanide or a 
 solution of potassium thiocyanate. Result ? This test is made to detect 
 the presence of which ion ? Equations ? 
 
 K 2 Cr 2 7 + 4 H 2 S0 4 -> K 2 S0 4 + Cr 2 (SO 4 ) 3 + 4 H 2 O + ( 3 O). 
 6FeS0 4 + 3 H 2 S0 4 + ( 3 O) 3 Fe 2 (SO 4 ) 3 + 3 H 2 O. 
 
 3. By potassium permanganate in the presence of an acid. Measure 
 accurately 15 cm. 3 of a recently prepared solution of ferrous sulphate 
 (ferrous-ammonium sulphate) into a small clean beaker, and add 5 cm. 3 
 of dilute sulphuric acid. Clamp clean burette into a vertical position 
 and fill it with a solution of potassium permanganate (end shelf). Allow 
 this latter solution to drop slowly into the ferrous salt. The pink color 
 of the permanganate immediately disappears on stirring with a glass 
 rod (do not remove the rod from the solution). The color continues to 
 be destroyed until all of the ferrous salt is completely oxidized to the 
 ferric state, when a drop of permanganate added in excess imparts a 
 faint pink color to the liquid. This indicates the end point] i.e., that the 
 reaction is ended. Record the number of cm. 3 of potassium permanga- 
 nate required to oxidize the ferrous salt. 
 
 By applying suitable tests to portions of this liquid, prove that the 
 oxidation has been complete; i.e., that the ferrous ion is not present 
 and that ferric ions are present. Make a record of tests employed and 
 results secured. 
 
 Repeat the experiment, using 20 cm. 3 of ferrous sulphate solution. 
 Equations ? 
 
 4. Chlorine water, bromine water, and potassium chlorate in the 
 presence of strong hydrochloric acid, are other oxidizing agents fre- 
 quently used to convert ferrous into ferric salts. 
 
 5. Add 5 cm. 3 of pure concentrated nitric acid to 50 cm. 3 of tap-water; 
 evaporate to 15-20 cm. 3 and test for the presence of the ferric ion (?). 
 
TRANSITION ELEMENTS. 363 
 
 COBALT, CO. 
 
 At. Wt. 59.0 Sp. Gr. 8.5. 
 
 The principal ores of cobalt are smaltite, CoAs 2 , and cobaltite, CoAsS. 
 The pure metal may be obtained by reducing the oxide, the chloride 
 or the oxalate in a stream of hydrogen or by Goldschmidt's process. 
 Cobalt resembles iron in many respects. It is a lustrous silver-white 
 (pink-tinted), hard metal, malleable, tenacious, and when heiated is very 
 ductile. It melts at about 1500. Unlike iron and nickel, t retains its 
 magnetic properties even at red heat. The metal has but few commer- 
 cial applications, its use being confined to the iron and steel industry. 
 
 Cobalt in the massive form is not readily acted upon by the air, but 
 the finely divided metal, especially when it is freshly prepared by the 
 reduction of the oxide in hydrogen, oxidizes easily and may take fire 
 spontaneously in the air. It is attacked slowly by dilute acids, hydrogen 
 being liberated. 
 
 Cobalt, like iron, forms two kinds of ions the cobaltous ion, Co", 
 and the cobaltic ion, Co'". The cobaltous salts are but slightly hydro- 
 lyzed, but the cobaltic salts are completely decomposed by water. The 
 latter are rather unstable and tend to break down into the cobaltous 
 salts with a liberation of one-third of the acid radical. Most of the 
 cobaltous compounds are red when hydrated and blue when dehydrated. 
 The blue color of the salt is explained by some chemists as being due to 
 the " repression of the ionization of the salt"; i.e., u to the driving back 
 of the ions into molecules, which are blue." Cobalt shows a marked 
 tendency to enter into combination with the ions, NO 2 , CN and NH 2 , 
 to form derivatives which yield many complex ions which usually give 
 none of the reactions of cobalt ions. The cobalt compounds give with 
 ammonia many complex compounds which present some very complex 
 relations. One of the most interesting series of these complex salts is 
 known as the cobalt amines. 
 
 Some of the more important compounds are represented by the fol- 
 lowing formulae: 
 
 Cobaltous compounds, CoO, Co(OH) 2 , CoCl 2 , 6H 2 O, CoSO 4 ,6H 2 O, 
 Co(NO 3 ) 2 .6H 2 O, CoCO 3 .6H 2 O, CoS, Co 2 (Fe(CN) 6 ), Co 3 (Fe(CN) 6 ), 
 Co(CN) 2 , K 4 Co(CN) 6 ; cobaltic compounds, Co 2 O 3 , Co(OH) 3 , 
 Co 2 (SO 4 ) 3 .i8H 2 O, K 3 Co(NO 2 ) 6 ; cobalt amines, Co(NH 3 ) 3 Cl 3 .H 2 6, 
 Co(NH 3 ) 4 Cl 3 .H 2 O, Co(NH 3 ) 4 Cl 3 .H 2 O, Co(NH 3 ) 5 Cl 3 , etc. 
 
 Experiment I. Properties of Cobalt. 
 
 Examine cobalt metal and note its most obvious physical properties ( ?) 
 Scratch the metal with the point of a knife blade. Is the metal hard? 
 Does it tarnish readily in the air? 
 
 Experiment II. Cobalt Compounds. 
 
 Examine the compounds of cobalt (end shelf). Give the color and 
 
364 EXPERIMENTAL CHEMISTRY. 
 
 empirical formula of each compound. Record the valency of the cobalt 
 atom in each compound. What is the color of the cobalt ion ? 
 
 Experiment III. Dehydration of Hydrated Cobaltous Chloride. 
 
 Write upon a sheet of your note-book with a solution of cobalt chloride 
 by means of a glass rod. Allow it to dry, then warm the paper very 
 gently by holding it at some distance from a gas flame. Result ? Breathe 
 upon the paper or hold it for an instant in a current of steam. Explain 
 the changes in color. Equations? What is " sympathetic ink "? 
 
 Experiment IV. Reactions of Cobalt Salts. 
 
 Treat separate portions of a solution of cobalt chloride with the fol- 
 lowing reagents: 
 
 (a) Test the solution with a borax bead in the oxidizing (?) and re- 
 ducing (?) flames. 
 
 (b) Ammonium hydroxide, first in small quantities (?), then in excess 
 (?). Equations? 
 
 (c) Sodium hydroxide solution, first in small quantities (?), then in 
 excess (?). Add bromine water and boil (?). Equations? 
 
 (d) Ammonium sulphide solution (?). Filter. Try the effect of 
 dilute hydrochloric acid upon the precipitate (?). Equations? 
 
 (e) Potassium cyanide solution, first in small quantities (?), then in 
 excess ( ?) ; add sodium hydroxide solution in considerable quantity, then 
 bromine water* until the color of bromine persists; warm gently. 
 Results (?) Equations? 
 
 NICKEL, Nl. 
 
 At. Wt. 58.7 Sp. Gr. 8.8-9.0 
 
 This element possesses many properties in common with cobalt. 
 It occurs chiefly in combination with arsenic as niccolite, NiAs, and 
 nickel glance, NiAsS. It is now manufactured chiefly from garnierite, 
 H 4 Ni 2 Mg 2 (SiO 4 ) 3 .4H 2 O, a silicate found in Australia. The crude 
 nickel may be obtained from granierite by reducing the ore in a blast 
 furnace or by electrolysis. Pure nickel is usually prepared by reducing 
 the oxide with carbon at a high temperature or by reducing the oxide in a 
 stream of hydrogen. 
 
 Nickel is a lustrous, white metal (yellow-tinted), very hard and tena- 
 cious, with a melting point at 1570. It takes a very high polish. It 
 tarnishes very slowly, even in moist air. On account of its resistance 
 to oxidation, it is used extensively as a protective covering for other 
 metals which are more readily oxidized, such as iron, etc. The process 
 of depositing one metal upon another by electrolysis is known as electro- 
 plating. Nickel forms a number of valuable alloys. German silver is an 
 alloy of copper, nickel and zinc (2 to i to i). Our co-called "nickel" of 
 
 * Bromine is much more solublein an aqueous solution of potassium bromide than 
 in water. 
 
TRANSITION ELEMENTS. 365 
 
 currency contains 75 per cent, copper and 25 per cent, nickel. Nickel- 
 steel contains from 4 to 15 per cent, of nickel. 
 
 Hydrochloric acid and sulphuric acid attack nickel with difficulty, 
 but nitric acid acts upon it very readily. Concentrated nitric acid renders 
 the metal "passive." 
 
 Nickel forms a bivalent ion, nickelion, Ni", which is of green color. 
 Nickel can form a higher stage of oxidation, for example, Ni 2 O 3 and 
 Ni(OH) 3 , but they are extremely unstable and do not behave as base- 
 forming compounds. 
 
 Nickel can also form complex ions, but these are neither so numer- 
 ous nor so stable as those of cobalt; this forms the essential difference 
 between the chemical conduct of these two elements. The complex 
 ions of nickel containing ammonia also differ from those of cobalt not 
 only in being derived from bivalent nickel, but also in being very unstable. 
 
 Most of the nickel salts are green when hydrated and yellow when 
 dehydrated. 
 
 The following formulae show the composition of some of the more 
 important compounds of nickel: Nickelous compounds, NiO, Ni(OH) 2 , 
 NiCl 2 .6H 2 O, NiSO 4 .6H 2 O, Ni(NO 3 ) 2 .6H 2 O, NiCO 3 .6H 2 O, NiS, 
 Ni 2 (Fe(CN) 6 ), Ni 3 (Fe(Cn) 6 ) 2 , Ni(CN) 2 , Ni(CN) 2 , 2 KCN; nickelic com- 
 pounds, Ni 2 O 3 , Ni(OH) 3 . 
 
 Experiment I. Properties of Nickel. 
 Same as Experiment I, " Cobalt." 
 Experiment II. Nickel Compounds. 
 Same as Experiment II, " Cobalt." 
 Experiment III. Reactions of Nickel Salts. 
 
 Same as Experiment IV, "Cobalt." A solution contains both the 
 nickel and cobalt ions. How may the cobalt be separated from the 
 nickel ? 
 
 PROBLEMS. 
 
 1. It was found that 1.586 grams of iron formed 2.265 grams of ferric 
 oxide. Calculate the atomic weight of iron. 
 
 2. How many grams of potassium bichromate will be required to 
 oxidize 10 grams of ferrous sulphate? 
 
 3. How many grams of potassium permanganate will be required to 
 oxidize 10 grams of ferrous sulphate? 
 
 4. What weight of stannous chloride will be required, theoretically, 
 to reduce i gram of ferric chloride to ferrous chloride ? 
 
CHAPTER XXXIV. 
 TRANSITION ELEMENTS 
 
 Second Long Period. Fourth Long Period. 
 
 (Ruthenium, Ru. 101.7) (Osmium, Os. 191.0) 
 
 (Rhodium, Rh. 103.0) (Iridium, Ir. 193.0) 
 
 (Palladium, Pd. 106.5) Platinum, Pt. 194.8 
 
 These rare metals constitute two separate transitional groups, yet 
 they are very closely related to one another. Again, they possess certain 
 properties which are markedly different. The first three have atomic 
 weights which are close to one hundred, while the atomic weights of the 
 last three are close to two hundred. The specific gravities show a similar 
 relation, as may be seen from the following table: 
 
 I. Ru,sp.gr. 12.26; Rh, sp. gr. 12.10; Pd, sp. gr. 11.9; 
 II. Os, sp. gr. 22.38; Ir, sp. gr. 22.4; Pt, sp. gr. 21.45. 
 
 All of these elements, however, resemble platinum more or less closely. 
 They are therefore spoken of as the platinum elements. In nature they 
 occur associated together in what is commonly known as platinum ore. 
 This ore, which is sometimes spoken of as native platinum, is found in 
 small particles and nuggets in river sand and alluvial deposits. The 
 Urals furnish the larger portion of the world's supply; smaller quantities 
 are found in Australia, California, Borneo and Brazil. The ore con- 
 tains these elements in the metallic state, more or less alloyed, together 
 with small quantities of iron, copper and gold. Platinum constitutes 60 
 to 85 per cent, of the ore. Smaller amounts of platinum are found alloyed 
 with iridium. 
 
 The members of these groups are lustrous, white metals, unacted upon 
 by air at ordinary temperatures. Osmium burns in oxygen when highly 
 heated, forming the tetroxide, OsO 4 . The other five metals resist oxi- 
 dation at any temperature. 
 
 Palladium is the only member which is attacked by nitric acid. The 
 other members are not acted upon by ordinary acids. Aqua regia is 
 without action upon rhodium and iridium. 
 
 Rhuthenium is a hard, brittle metal, fusing at about 2000. It was 
 discovered by Claus in 1845. Some of its compounds are RuO, Ru 2 O 3 , 
 RuO 2 , RuO 4 , Ru(OH) 3 , Ru(OH) 4 . 3 H 2 O, RuCl 2 , RuCl 3 , RuCl 4 , K 2 RuO 4 , 
 KRoO 4 . 
 
 Rhodium is a malleable metal, fusing at 2000. In appearance, it re- 
 sembles aluminum. It is harder than platinum. It was discovered by 
 
 366 
 
TRANSITION ELEMENTS. 367 
 
 Wollaston in 1803. Some of its compounds are, RhO, Rh 2 O 3 , RhO 2 , 
 Rh(OH) 3 , Rh(OH) 4 , RhCl 3 . 
 
 Palladium is the most easily fusible of these metals, melting at about 
 1500. It possesses the property of absorbing (occluding) large volumes 
 of hydrogen gas, 360 to 960 times its own volume, depending upon its 
 state of aggregation and its temperature. When heated to 130, it 
 surrenders the hydrogen. It has been supposed for a long time that the 
 hydrogen and palladium enter into a definite chemical union with the 
 formation of palladium hydride, Pd 2 H. There is considerable doubt as 
 to whether this is a definite chemical compound. The later explanation 
 that palladium-hydrogen is simply a "solid solution" in which hydrogen 
 is dissolved in palladium does not satisfactorily account for all of the 
 observed phenomena. Palladium-hydrogen is a powerful reducing agent 
 owing to its capability of releasing hyrogen in a condition similar to 
 "nascent" hydrogen Some of its compounds are Pd 2 O, PdO, PdCl 2 , 
 PdCl 4 , PdI 2 , Pd(NO 3 ) 2 , PdSO 4 .2H 2 O, H 2 PdCl 6 , K 2 PdCl 6 , (NH 4 ) 2 PdCl 6 . 
 
 Osmium is the heaviest known elementary substance. It fuses at 2 500 . 
 An alloy of iridium and osmium which is very hard is used for tipping gold 
 pens. A solution of osmium tetroxide, OsO 4 , is found useful in hardening 
 tissues for histological purposes. Some of its more important compounds 
 are represented by the following formula?: OsO, Os 2 O 3 , OsO 2 , OsO 4 , 
 Os(OH) 2 , Os(OH) 3 , OsCl 2 , OsCl 3 , OsCl 4 , K 2 OsO 4 . 
 
 Iridium is a very hard metal fusing at 1950. It alloys with platinum 
 and enhances the resistance of that metal to the action of acids. It is, 
 therefore, usually present in platinum utensils designed for laboratory 
 purposes. The following formulae represent some of its compounds: 
 IrO, Ir 2 3 , Ir0 2 , Ir(OH) 4 , IrCl 2 , IrCl 3 , IrCl 4 , K 3 IrCl 6 . 3 H 2 O. 
 
 Platinum is a tough, malleable, ductile metal which can be welded at 
 red heat It fuses at 1770. Its temperature coefficient of expansion is 
 about the same as that of glass. Therefore, if it is desired to seal an elec- 
 trical connection through glass, a platinum wire which is a good con- 
 ductor is the most convenient means. This fact is utilized in construct- 
 ing incandescent electric lights. 
 
 Platinum occludes oxygen and hydrogen, the quantity absorbed de- 
 pending upon the state of division of the metal. Platinum black (very 
 finely divided platinum) which can be prepared by depositiong it from 
 its solutions by means of a more electro-positive metal, absorbs* about 
 300 times its own volume of hydrogen and about 100 times its own 
 volume of oxygen. At red heat the gases are expelled from the metal. 
 The value of platinum as a "catalytic agent" is due to its capacity to 
 occlude gases. 
 
 Platinum is very resistant to chemical reagents, and upon this fact its 
 value largely depends. It is not attacked by the ordinary acids, but the 
 free chlorine in aqua regia converts it into chloroplatinic. acid, H 2 PtCl 6 . 
 It is not acted upon by the fused alkaline carbonates, but it interacts with 
 the fused alkalies, giving platinates. The, fused alkaline cyanides also 
 
 * Hydrogen is "dissolved," but oxygen is merely concentrated upon its surface. 
 
368 EXPERIMENTAL CHEMISTRY. 
 
 interact with it. Platinum must not be heated in contact with carbon, sili- 
 con and phosphorus, as they unite with it forming compounds which are 
 quite brittle. Lead and antimony form fusible alloys with platinum, 
 therefore neither these metals, nor compounds from which they may be 
 liberated, should be heated in platinum vessels. 
 
 Platinum forms compounds in which it is the positive ion showing a 
 valence of II or IV, as well as compounds in which it is a constitutent of 
 the negative ion. However, when solutions containing ions, in which 
 platinum is in the anion, are electrolyzed, the platinum is deposited at the 
 anode and not at the cathode. The following list contains some of the 
 more important compounds of platinum: Platinous compounds, PtO, 
 Pt(OH) 2 , PtCl 2 , PtS; Platinic compounds, Pt(OH) 4 , PtCl 4 , PtS 2 ; chloro- 
 platinic acid, H 2 PtCl 6 ; chloroplatinates, K 2 PtCl 6 , Na 2 PtCl 6 , (NH 4 ) 2 PtCl 6 ; 
 platinocyanides, BaPt(CN) 4 .4H 2 O, K 2 Pt(CN) 4 ,3H 2 O. 
 
 Experiment I. Properties of Platinum. 
 
 (a) Physical properties. Examine specimens of platinum foil and 
 wire. Note the physical properties of the metal (?). Hold a piece of 
 platinum wire in the hottest portions of the Bunsen flame (?), and the 
 flame of the blast-lamp (?). What is "spongy platinum"? " Platinum 
 black"? 
 
 (b) Chemical properties. Does platinum tarnish readily when ex- 
 posed to the action of the air. Procure two small pieces of platinum 
 scrap from the assistant. Place them in separate test tubes. Heat the 
 one with hydrochloric acid (?) and the other with nitric acid (?). Mix 
 the contents of both test tubes. Result ? Equation ? 
 
 Experiment II. Platinum as a " Catalytic Agent." 
 
 Recall or repeat those experiments in which platinum acted -as a 
 " catalytic agent." State briefly Ostwald's tentative explanation of the 
 role of a catalyzer. (See Chapter IX, Experiment VII.) 
 
 Experiment III. Platinum Compounds. 
 
 Examine the compounds of platinum (end shelf). Give the color and 
 formula of each compound. What is the color of the platinum ion? 
 
 Experiment IV. Preparation of Platinic Chloride. 
 
 (a) Scour pieces of platinum scrap with sea sand, then wash with dis- 
 tilled water, and boil with hydrochloric acid. Decant the liquid, and 
 again wash with distilled w r ater. Dissolve i gram of the platinum in 25 
 cm. 3 of aqua regia in a covered glass dish. After the "spurting" has 
 ceased pour off the supernatant fluid and concentrate it in a glass dish by 
 evaporation on the steam bath. In the meantime add more aqua regia 
 to the undissolved platinum, and continue the above operations until all 
 of the platinum is dissolved. After the combined solutions have been 
 evaporated to dryness, moisten with a little hydrochloric acid and take 
 up with water, or, 
 
TRANSITION ELEMENTS. 369 
 
 (b) The platinum may be precipitated from the concentrated aqua- 
 regia solution by a strong ammonium chloride solution, as ammonium 
 chloroplatinate, (NH 4 ) 2 PtCl 6 . Filter, ignite the precipitate in a porcelain 
 crucible. Boil the residue of "spongy platinum" with hydrochloric acid, 
 decant the liquid and dissolve the platinum in aqua regia. Evaporate the 
 solution to dryness on the steam bath, moisten the residue with hydro- 
 chloric acid and again evaporate just to dryness. Take up the product 
 with distilled water. Equations? 
 
 Experiment V. Reactions of Platinum Salts. 
 
 (a) Dilute a small quantity of a platinic chloride solution with 2 or 3 
 cm. 3 of water, add a solution of hydrogen sulphide, and warm slightly ( ?). 
 Is the precipitate soluble in excess of potassium hydroxide ? In ammo- 
 nium polysulphide ? 
 
 (b) Reduction of platinum salts. "Platinum black".* To a solution 
 of platinic chloride add an excess of sodium carbonate solution and a little 
 grape sugar, then boil the mixture. Carbon dioxide is evolved, and a 
 black powder ("platinum black") is formed slowly. The solutions 
 should be dilute. Wash the powder successively with dilute alcohol, 
 hydrochloric acid, potassium hydroxide solution, distilled water, and then 
 dry it by applying a gentle heat. 
 
 (c) Recall the interactions of solutions of platinic chloride with the 
 alkaline chlorides. (See "Sodium," "Potassium," and "Ammonium.") 
 
 For a discussion of the properties of praseodymium, neodymium, sama- 
 rium, gadolinium, terbium, erbium, thulium, ytterbium, etc., see 
 reference texts. 
 
 *" Platinum black" is a more effective catalyzer than "spongy platinum." 
 
CHAPTER XXXV. 
 
 RELATIONS WITHIN THE GROUPS or THE PERIODIC CLASSIFICATION. 
 
 TABLE. 
 Grouping of the Metals (Cations) for Purposes of Analysis. 
 
 METALS (CATIONS) 
 
 _ ! __ 
 
 II. 
 
 Div. A. 
 
 Div. B. 
 
 III. 
 
 IV. 
 
 V. 
 
 Metals which are precipitated 
 from solutions by 
 
 HC1 
 Silver 
 
 Mercury (ous) 
 Lead 
 
 I 
 
 Metals which are not precipitated 
 from solutions by 
 
 HC1 
 
 Sulphides soluble in ammonium polysulphide. 
 
 Metals which are precipitated 
 from solutions containing 
 hydrochloric acid by 
 H 2 S 
 
 Mercury (ic) 
 
 Copper 
 
 Cadmium 
 
 Bismuth 
 
 Arsenic 
 
 Antimony 
 
 Tin 
 
 Gold 
 
 Platinum 
 
 Metals precipitated from ammoni- 
 acal solutions containing ammo- 
 nium chloride by 
 (NH 4 ) 2 S 
 
 Cobalt 
 
 Nickel 
 
 Iron Chromium 
 
 Aluminum 
 
 Zinc 
 
 Manganese 
 
 I 
 
 Metals which are not precipitated 
 from solutions containing hy- 
 drochloric acid by 
 
 H 2 S 
 
 Metals not precipitated from ammoni- 
 acal solutions containing ammonium 
 chloride by 
 
 (NH 4 ) 2 S 
 
 I 
 
 Metals which are precipitated 
 from solutions by 
 
 (NH 4 ) 2 C0 3 
 Barium 
 Calcium 
 Strontium 
 
 Metals which are not precipitated 
 from solutions by 
 
 (NH 4 ) 2 C0 3 
 
 Metal which is precipitated 
 from solutions by 
 NaNH 4 HPO 4 
 
 Magnesium 
 
 VI. 
 
 Metals which are not precipitated 
 from solutions by 
 
 NaNH 4 HPO 4 
 Sodium 
 Potassium 
 Ammonium 
 
 370 
 
EXPERIMENTAL CHEMISTRY. 
 
 371 
 
 TABLE. 
 
 GROUPING or NON-METALLIC RADICALS (ANIONS) FOR PURPOSES or ANALYSIS. 
 
 ACID-RADICALS (ANIONS) 
 
 I 
 
 Radicals which are precipitated 
 from neutral solutions bv 
 
 BaCl 2 
 
 J 
 
 Radicals which are not precipitated 
 from neutral solutions by 
 
 BaCl 2 
 
 Sulphuric, H 3 SO 4 
 Hydrofluosilic, H 2 SiF 6 
 
 Boric, H 3 BO 3 
 Carbonic, H 2 CO 3 
 Citric, H 3 (C 6 H 5 6 7 ) 
 Chromic, H 2 CrO 4 
 Hydrofluoric, HF 
 lodic, HIO 3 
 Oxalic, H 2 (C 2 O 4 ) 
 Phosphoric, H 3 PO 4 
 Sulphurous, H 2 SO 3 
 Silicic, H 2 SiO 3 
 Tartaric, H 2 (C 4 H 4 O 6 ) 
 Thiosulphuric, H 2 S 2 O 3 
 
 Barium salts insoluble in dilute 
 hydrochloric acid. 
 
 Barium salts soluble in dilute 
 hydrochloric acid. 
 
 I 
 
 II. Radicals precipitated from solutions 
 acidified with nitric acid by 
 
 AgN0 3 
 
 Radicals which are not precipitated 
 from solutions acidified with nitric 
 acid by 
 
 AgN0 3 
 
 Ferricyanic, H 3 Fe(CN) 6 
 Ferrocyanic, H 4 Fe(CN)6 
 Hydrochloric, HC1 
 Hydrobromic, HBr 
 Hydriodic, HI 
 Hydrocyanic, HCN 
 Hydrosulphuric, H 2 S 
 Thiocyanic, HCNS 
 
 Acetic, H(C 2 H 3 O 2 ) 
 Chloric, HC1O 3 
 Perchloric, HC1O 4 
 Cyanic, HCNO 
 Formic, H(CHO 2 ) 
 Nitric, HNO 3 
 Nitrous, HNO 2 
 
APPENDIX I. 
 
 PRELIMINARY EXERCISES. 
 APPARATUS. 
 
 Since chemistry is pre-eminently an experimental science, it is in order 
 to begin work by examining various pieces of apparatus and acquainting 
 one's self with their manipulation. 
 
 THE BUNSEN BURNER. 
 
 Carefully examine the Bunsen burner (Fig. 49). Take it apart; assemble 
 it, noting relation of parts. Connect it with the gas supply. Light the 
 burner by turning on the gas, and then holding a lighted match near the 
 side of the burner a short distance below the top. Notice the effect of 
 opening and closing the holes at the bottom of the tube. Explain. Ad- 
 just the opening until the luminous region has 
 just disappeared and the flame is noiseless: 
 this is the "Bunsen flame." A flame 6 cm. to 
 9 cm. high is adapted to most work. If soot is 
 deposited upon the object being heated, open 
 the holes at the foot at the tube, but not so far 
 as to produce a noisy flame. 
 
 Now study the structure of the flame. De- 
 termine which parts are relatively hotter and 
 which cooler. Introduce quickly into the 
 center of the flame (the darker portion), about 
 one-half of a centimeter above the burner, the 
 head of a match. Results ? Can you insert a 
 match as per above directions and withdraw it 
 without igniting same? 
 
 Hold a piece of platinum wire across the 
 flame in various places. Note the color of the 
 wire. Bring quickly and horizontally (to desk 
 top) a piece of heavy pasteboard into the flame, 
 about 2 cm. above the top of the tube; hold it quietly for several seconds, 
 then remove quickly. Do you find a brown (scorched) circle ? Explain. 
 
 To secure the greatest heating effect, place object just above apex of 
 the dark inner cone of unburned gas. Why is there a cone of unburned 
 gas ? Why the shape ? If oxygen is necessary for the combustion of the 
 gas, where might the gas escape for a time unburned? 
 
 373 
 
 FIG. 49. Bunsen Burner. 
 
374 EXPERIMENTAL CHEMISTRY. 
 
 Half fill a test tube with water. Be sure that the outside of the tube is 
 dry. By means of test tube holders, introduce the tube into the hottest 
 region of the flame, inclining the tube at an angle of about 45 to the top 
 of the desk. Heat only that portion of the tube containing the liquid: 
 if the flame strikes the tube above the liquid, the tube may crack. 
 Keep the liquid in the tube slightly agitated by a short, quick move- 
 ment of the hand. 
 
 Above directions should be observed in all cases where liquids are 
 heated in test tubes. A "wing-top" attachment gives a broad flame 
 of very much use to the glass-blower. 
 
 Note. If gas is not available in the laboratory and alcohol lamps are 
 used, perform as many as possible of above experiments. 
 
 THE BLAST-LAMP. 
 
 Where a much higher temperature is required than can be secured by 
 means of the Bunsen burner, the blast-lamp is used. The size of the 
 flame can be altered by proper manipulation. (Instructions from assistant.) 
 
 MANIPULATION OF GLASS. 
 
 I. To Cut Glass Tubing. 
 
 (a) Lay the tubing on a flat surface; make a file-mark on it at right 
 angles to the length; take the tube in the hands, placing the two thumbs 
 opposite the scratch and the fingers on either side of the scratch; now push 
 gently with the thumbs and at the same time pull the hands apart, the 
 tubing usually breaks squarely at the scratch. " Fire-polish" the ends 
 of tube by turning them slowly in the Bunsen flame. 
 
 (b) To break large tubing or cut off bottoms of bottles, etc., encircle 
 tube with wire or make an ink-mark to trace the path of the desired 
 break or cut, then a file-mark is made upon the surface on mark; a steel 
 file handle, or better, a glass rod heated in the blast-lamp flame until it is 
 red hot, when it is at once pressed against the scratch until the glass 
 begins to crack. The fracture can usually be led in any direction by keep- 
 ing the hot glass rod in front of it. Heat rod frequently to keep it red hot. 
 
 II. Grinding Glass. 
 
 (a) Rough edges of tubes or bottles may be ground down to bell-jar 
 effect by spreading emery paste upon a smooth flat surface and rubbing 
 broken edges upon it. 
 
 (b) Glass stoppers may be ground into necks of flasks, etc., by covering 
 stoppers and inside of neck with emery paste, and then imparting a gentle 
 pressure to stopper while twisting it into place. 
 
 III. Cutting and Perforation Glass Plates. 
 
 (a) The plate of glass is laid upon a flat surface; "glass cutters" 
 are used to make the scratch where the break is desired. " Glass cutters" 
 are usually made with a rotating wheel of steel or a diamond point. 
 
APPENDIX I. 375 
 
 (b) Holes can be made in a glass plate by the aid of a broken end of a 
 round file kept wet with a solution of camphor in oil of turpentine. 
 
 IV. To make Stirring Rods. 
 
 Cut off a piece of glass tubing 18 cm. to 20 cm. long and 6 mm. in di- 
 ameter. Hold the ends of the tube successively in a Bunsen flame, rotating 
 the tube constantly until the open ends are sealed. Glass rods may be 
 used instead. 
 
 V. To Bend Glass Tubing. 
 
 A flat Bunsen flame, produced by a "wing-top" or a "fish-tail" attach- 
 ment, is used for bending glass. Take the tube in both hands and hold 
 that portion which is to be bent lengthwise to the flame and just above the 
 flame until it is warmed; then place it in the flame, constantly rotating on 
 the long axis, between thumb and fingers, until glass becomes fairly soft; 
 remove it from flame and quickly bend it into the desired form. It is 
 well to anneal the glass at the bend by "smoking" it. This maybe ac- 
 complished by closing the holes at the base of the burner, thus producing 
 a smoky flame. The bent portion should not be permitted to touch cold 
 objects until it has cooled. 
 
 Using ordinary glass tubing, the student should make various styles of 
 bends like models shown by the assistant. Always "fire-polish" the 
 edges of glass tubing. 
 
 Note. To increase the internal diameter of tubes for insertion of corks, 
 etc., soften the tube in the flame, and insert a conical piece of charcoal 
 by gentle pressure until tube spreads into desired shape. 
 
 VI. Joining Tubes and Glass Blowing. 
 
 See Ostwald's "Physico-Chemical Measurements," pages 66-72, also 
 Shenstone's " Methods of Glass Blowing." The student will find in said 
 references brief but excellent discussions of the subject. 
 
 VII. Sealing Platinum Wires into Glass Tubes. 
 
 Soften a glass tube; draw it out; cut off short, and by heating cause the 
 end to fall nearly together, or until the wire can just be pushed into the 
 opening; heat until glass closes around the wire. 
 
 More certain results are obtained if the platinum wire receives a drop 
 of melted enamel (tough lead-glass) at the proper place and is pushed 
 through opening in tube until hole is closed by the enamel. The enamel 
 unites well with platinum and also with ordinary glass. 
 
 PERFORATION OF STOPPERS. 
 
 A set of cork borers may be secured from the assistant. (T.O.) 
 Hold the cork in the hand and bore from the narrow end. Avoid great 
 pressure on the cork borer. 
 
376 
 
 EXPERIMENTAL CHEMISTRY. 
 
 In perforating rubber corks, the borer cuts more easily if it is dipped 
 frequently into a solution of caustic soda. A round file may be used to 
 smooth the perforation. 
 
 TREATMENT OF RUBBER CORKS AND TUBING. 
 
 Where rubber is to be used in quantitative experiments, it should be 
 boiled in dilute sodium hydroxide solution, rinsed with water, then boiled in 
 dilute hydrochloric acid, and finally washed with water. This operation 
 removes impurities which frequently introduce errors. 
 
 FIG. 50. Water Bottle. 
 
 FIG. 51. Mohr Burettes. 
 
 CONSTRUCTION OF PARTS AND THE ASSEMBLING OF A WASH BOTTLE. 
 
 Select proper material and construct a wash bottle (Fig. 50) like model 
 in laboratory. The necessary material will be found in the drawer. 
 
 MEASURING INSTRUMENTS. 
 
 (a) Measures of volume. 
 
 Assemble the Mohr (Fig. 51) ; or Geissler burette, and clamp it in a verti- 
 cal position. Observe the model. Fill the burette with distilled water ; avoid 
 air bubbles in stop cock; run out water into a beaker until the lower side 
 of meniscus (Fig. 54) (curved surface) stands at 40.85 c.c.; estimate to 
 
APPENDIX I. 
 
 377 
 
 tenths of a division; record this reading in your note-book; now run out 
 water into a "graduate " until the latter is about half filled; take the read- 
 ing on the "graduate"; how many cm.*? Take the burette reading; 
 how many cm. 3 did you run out ? Compare readings. Which gives the 
 more accurate measurement burette or graduate ? Why ? 
 
 By means of a calibrated pipette (Fig. 55), introduce its nominal volume 
 of water into a "graduate." Compare readings. Which is the more 
 accurate ? Why ? 
 
 By use of the graduate, determine the approximate volumes of water 
 which test tubes, beakers and flasks will hold. Make a record of data. 
 
 FIG. 52. Burette 
 
 Operated with 
 
 Pinch-cock. 
 
 FIG. 53. Burettes with Glass 
 Stop-cocks. 
 
 FIG. 54. Meniscus. 
 
 By aid of burette, determine the volume occupied by a "drop" of water 
 as delivered by your burette. Hint. Run out ten to twenty drops. Make 
 a record of exercise. 
 
 (b) Measure of temperature. 
 
 The mercury-glass therometer is ordinarily used in the laboratory. 
 Note the possible maximum and minimum readings. Always observe 
 this before using instrument. Some thermometers are made to read to 
 only -f 60 C.; such a thermometer plunged into water boiling at any 
 temperature above that registered on scale would undoubtedly be ruined. 
 Never plunge the instrument into relative extremes of temperature. 
 
 Suspend the thermometer in a beaker of water which has presumably 
 acquired the temperature of the laboratory. Remember the heat of the 
 
EXPERIMENTAL CHEMISTRY. 
 
 25C 
 
 hand will affect the reading. It requires time for a thermometer to come 
 to the temperature of a new environment the thermometer is said to 
 "lag." Gently tap the thermometer with finger before reading 
 this is to overcome "stiction." What is the temperature of the water in 
 beaker on the Centigrade scale? Fahrenheit scale? Reaumur scale? 
 Make a record of experiment. 
 
 The Beckmann thermometer for deter- 
 mining accurately small changes in temper- 
 ature is frequently used. Consult the in- 
 structor with reference to its manipulation. 
 (c) Measures of weight. 
 The equal arm lever balance is the in- 
 strument most frequently used in the 
 chemical laboratory for the determination 
 of weight. The principle of this balance is 
 embodied in two useful forms the "plat- 
 form" or "trip" balance and the so-called 
 "analytical" balances (Fig. 56). Both 
 kinds will be found in the laboratory. Ex- 
 amine them, the latter under the supervision 
 of the assistant. The former is used in 
 making those weighings where only an ap- 
 proximate accuracy is demanded; the latter 
 are much more sensitive and accurate, and 
 should be used only when the experiment 
 is marked "Quant." 
 
 The balance and weights must be handled 
 with care. Solids to be weighed must be 
 placed first upon a piece of paper or a watch 
 glass, never directly upon the pan. All 
 objects must be perfectly clean and dry. 
 The weights must be handled with forceps 
 not with the moist hand. 
 
 The general procedure when weighing 
 with platform balances is as follows: find 
 zero point of balance by allowing the beam 
 to swing, noting whether the pointer makes 
 equal excursions on either side of the zeVo 
 mark; if it does not, correct defects by 
 placing pieces of paper, etc., on the proper 
 pan. This is called "counterpoising" a 
 
 -pic. 55. Pipettes. balance. In the future, when weighing, do 
 
 not wait for pointer to come to rest at the 
 
 zero point, simply add to or subtract weights from the proper pan until 
 the vibrations on either side of zero point are of equal amplitude. 
 
 The assistant will instruct you as to the proper method of weighing with 
 the "analytical" balances. 
 
APPENDIX I. 
 
 379 
 
 Note. To avoid one of the commonest errors in weighing, count the 
 values of the vacant places in the set, then check by counting the weights 
 in the pan. Record on paper the value of the weights as you take them 
 from the pan. 
 
 CALIBRATING BY WEIGHING. 
 
 Make a mark or paste a piece of gummed paper on the lower portion 
 of the neck of a 50 cm. 3 Erlenmeyer flask. Weigh it first on the platform 
 balances, then on the analytical balances. (Instructions.) Place flask 
 on left pan. Record weights. Now fill flask to mark with water at 
 
 
 
 
 
 FIG. 56. Chemical Balance. 
 
 20 C. and weigh as before. Record weights. If a cm. 3 of water at 
 20 C. weighs .998 gram, what is the volume of flask when filled to 
 mark? Which is buoyed up the most by air flask and contents or 
 weights? Is this a source of error? 
 
 PROBLEMS. 
 
 1. Express 10 cm. in millimeters, decimeters and kilometers. 
 
 2. Convert 115 in. into centimeters; 98 mm. into inches; 2 yd. into 
 meters; 2 ft. into centimeters; 760 mm. into inches. 
 
 3. How many liters in 10 quarts? Quarts in 4 liters? 
 
 4. How many liters will a pneumatic trough hold, the dimensions of 
 which are 15 in.x2o in.x25 in.? 
 
 5. Add 9 grm., 468 mg., 7 dg., and 5 eg., and express the sum in grams. 
 
 6. 10 kg. are equivalent to how many pounds? Convert 10 oz. into 
 grams. 
 
380 EXPERIMENTAL CHEMISTRY. 
 
 7. How many grams in 134.76 dg. ? In 17,589 eg.? In 5.95 mg. ? 
 
 8. How many cubic centimeters (cm. 3 ) in .5!.? In .75!.? In 95 dm. 3 
 
 9. Convert 32 Fahr. into a Centigrade reading. A Reaumur read- 
 ing. 
 
 10. Repeat 9, using 212 Fahr. 
 
 11. A thermometer bearing a Fahrenheit scale registers a temperature 
 of 72. What would be the equivalent reading on the Centigrade scale? 
 On the Reaumur Scale? 
 
 12. Convert 20 C. into an equivalent reading on the Fahr. scale. 
 On the Reaumur scale. 
 
 13. The sp. gr. of concentrated sulphuric acid is 1.84, of nitric acid 
 1.4, and of hydrochloric acid 1.2. Calculate the weight of a liter of each 
 acid. 
 
 14. How many cm. 3 in 10 grm. of each of the above acids? 
 
 15. Alcohol (ethyl) at 15 C. has a density of .7937. What is the 
 weight of one 1. at 15 C. How many cm. 3 in 20 grm.? 
 
 1 6. An empty flask has a weight of 96.75 grm.; when filled with water 
 at 20 C. it has a weight of 596.30 grm. What is the capacity of the 
 flask at 20 C. ? Hint. calculate the number of grams of water which 
 the flask holds and multiply by the "correction factor," 1.0028. This 
 factor corrects for the buoyant effect of the air in weighing, the density of 
 water (.99823) at 20 C., and the cubical coefficient of expansion of glass 
 (.000025). The "correction factor" for 19 C. is 1.0027. See Appendix. 
 In standardizing volumetric apparatus the "correction factor" should 
 always be used. 
 
APPENDIX II. 
 
 NOTE. The data contained in the following tables have been selected 
 and arranged from various sources. Effort has been made to incorporate 
 the results of the most recent investigations. In many instances it has 
 been possible to give only approximate values. 
 
 The tables which have been most frequently consulted are: Landolt 
 und Bornstein, Physikalisch-chemische Tabellen; Ostwald, Physico- 
 chemical Measurements; Kohlrausch, Praktische Physik; Buchka, 
 Physikalisch-chemische Tabellen der unorganischen Chemie; Watson, 
 Practical Physics; Miller, Laboratory Physics; Smithsonian Physical 
 Tables; Comey, Dictionary of Solubilities; Seidell, Solubilities of Inorga- 
 nic and Organic Substances; J. Thomsen, Thermo-chemische Unter- 
 suchungen. 
 
 TABLE I. 
 Metric Measures with English Equivalents. 
 
 (a) Measures of Length. 
 
 i Millimeter, mm. = .0393 inches, 
 
 i Centimeter, cm. =10 mm. = .3937 inches, 
 
 i Decimeter, dm. =10 cm. = 3.9371 inches, 
 
 i Meter, m. =10 dm. =39.3708 inches, 
 
 i Meter, m. = 3 . 2809 feet, 
 
 i Meter, m. = 1.0936 yards, 
 
 i Kilometer, Km. =1000 m. = .6114 miles. 
 
 (b) Measures of Volume. 
 
 i Cubic centimeter, cm. 3 = .001 1. = .06103 cu - inches. 
 
 i Cubic decimeter (liter), i. = 1000 cm. 3 = 61.227 cu - inches. 
 
 i Liter = 2.1134 pints U.S. = 1.067 quarts U. S.= 0.26417 gallons U. S. 
 
 i Cubic meter = 1000 1. = 35.317 cu. ft. 
 
 (c) Measures of Weight. 
 
 Milligram, mg. .001 gram = .0154 grains. 
 
 Centrigram, eg. = .01 gram = .1543 grains. 
 
 Decigram, dg. = .1 gram = 1.5432 grains. 
 
 Gram, gm. = 15.432 grains = .0352 av. ounces. 
 
 Gram = .03216 troy ounces. 
 
 Kilogram (kilo) = 1000 grams = 2 . 2055 av. Ibs. 
 
 Kilogram == 2 . 6803 troy Ibs. 
 
 38* 
 
2 EXPERIMENTAL CHEMISTRY. 
 
 TABLE II. 
 English Measures with Metric Equivalents. 
 
 (a) Measures of Length. 
 
 i Inch =25.399 mm - = 2 -539 centimeters. 
 
 i Foot =12 inches = .3048 meter. 
 
 i Yard =3 feet = .9144 meter. 
 
 i Mile =1700 yards = 5280 feet. 
 
 i Mile = i . 609 kilometers = 1609.3 meters. 
 
 (b) Area. (c) Contents. 
 
 i Square inch = 6.4514 cm. 2 i Cubic inch = 16.386 cm. 3 
 
 i Square foot = 929.01 cm. 2 i Cubic foot = 28.316 liters, 
 
 i Square yard = 8361.1 cm. 2 i Cubic yard- = 764.52 liters, 
 
 i Square yard = .83661 m. 2 i Cubic yard = about .76 m. 3 
 
 (d) Measures oj Volume. 
 
 i Pint = .47317 liter = 473.11 cubic centimeters, 
 
 i Quart = .94634 liter = 946.22 cubic centimeters, 
 
 i Gallon(U.S.) = 3.785 liters = 231 cubic inches. 
 
 (e) Apothecaries' Fluid Measure, U.S. 
 
 i Minim = a drop (approx) .0616 cubic centimeters, 
 
 i Fluid dram = 60 minims 3 . 6965 cubic centimeters. 
 
 i Fluid ounce = 8 fluid drams = 29.572 cubic centimeters, 
 i Pint = 16 fluid ounces = 473.11 cubic centimeters. 
 
 (/) Measures of Weight. 
 
 I. (Avoirdupois.) 
 
 i Gram = 64.773 milligrams = 0.0648 grams, 
 i Dram = 27.34 grains = 1.772 grams, 
 i Ounce = 1 6 drams = 437.5 grains = 28.349 grams. 
 i Pound = 16 ounces = 7000 grains = 453.59 grams. 
 i Short ton = 2000 pounds = 907.17 kilograms, 
 i Long ton = 2240 pounds = 1015.03 kilograms. 
 
 II. (Troy.) 
 
 i Grain = 64.773 milligrams = .0648 grams. 
 i Pennyweight = 24 grains = 1.555 grams, 
 i Ounce = 20 pennyweights = 480 grains = 31.103 grams, 
 i Pound = 12 ounces = 5760 grains = 373.242 grams. 
 
 III. (Pharmacy or Apothecaries\) 
 i Grain = 64.773 milligrams = .0648 grams, 
 i Scruple = 20 grains = i . 296 grams, 
 i Dram = 3 scruples = 60 grains = 3 . 888 grams, 
 i Ounce = 8 drams = 480 grains = 31 . 1035 grams, 
 i Pound = 12 ounces = 373.248 grams. 
 
APPENDIX II. 
 
 TABLE III. 
 
 383 
 
 Conversion of Thermometric Readings. 
 
 The temperatures mentioned in this book are expressed in terms of 
 the Centigrade scale. There are three scales now in general use. They 
 
 are: 
 
 I. Fahrenheit F. Water freezes at 32, boils at 212. 
 II. Centigrade C. Water freezes at o, boils at 100. 
 III. Reaumur R. Water freezes at o, boils at 80. 
 
 To convert: F. to C. 
 
 = C. 
 
 5(F. 32) 
 or = C. 
 
 F. to R. 4 ( F -~32) = R> 
 
 C. to F. (C. x 1.8) + 32 = F.. 
 
 or 
 
 = F. 
 
 R.toF. 
 
 TABLE IV. 
 
 Corrections to Reduce Readings on Mercury-in-Glass Thermometers 
 to the Normal Hydrogen Scale or Air Thermometer. 
 
 (For Jena Normal Glass, 16'")* 
 
 Reading 
 Correction 
 
 oC. 
 o.ooo 
 
 10 C. 
 
 -55 
 
 20 C. 
 
 0.090 
 
 30 c. 
 
 0.109 
 
 40 C. 
 0.116 
 
 0.109 
 
 Reading 
 Correction 
 
 50 c. 
 
 o.io9 
 
 60 C. 
 
 0.096 
 
 70 C. 
 0.076 
 
 80 C. 
 -53 
 
 90 C. 
 
 - 0.027 
 
 100 C. 
 
 0.000 
 
 *From tables published by Griitzmacher in Wied. Annalen (1899) p. 769. 
 
EXPERIMENTAL CHEMISTRY. 
 
 TABLE V. DENSITY OF WATER.* 
 
 (Ax THE TEMPERATURE "/" ON THE NORMAL HYDROGEN SCALE) 
 
 8 
 
 1^ 
 
 Q 
 
 Tenths. 
 
 o 
 
 i 
 
 2 
 
 3 
 
 4 
 
 5 
 
 6 7 
 
 8 
 
 9 
 
 
 
 0-99987 
 
 87 
 
 88 
 
 89 
 
 89 
 
 90 
 
 90 91 92 
 
 92 
 
 i 
 
 2 
 
 3 
 4 
 
 93 
 97 
 99 
 
 I ' 00000 
 
 93 
 97 
 
 99 
 
 oo 
 
 94 
 
 97 
 99 
 
 00 
 
 94 
 
 98 
 
 00 
 00 
 
 95 
 98 
 
 00 
 00 
 
 95 
 99 
 
 00 
 
 oo 
 
 95 9 6 
 99 99 
 
 00 00 
 
 oo ; oo 
 
 96 
 99 
 
 00 
 00 
 
 97 
 99 
 
 00 
 
 00 
 
 5 
 
 0-99999 
 
 99 
 
 99 
 
 99 9 8 
 
 98 
 
 98 98 97 97 
 
 6 
 
 8 
 9 
 
 97 
 93 
 88 
 81 
 
 97 
 93 
 87 
 80 
 
 96 
 92 
 86 
 79 
 
 96 
 92 
 86 
 
 79 
 
 95 
 9 1 
 
 78 
 
 95 
 90 
 84 
 77 
 
 95 94 
 90 89 
 84 83 
 76 75 
 
 94 
 89 
 82 
 
 74 
 
 93 
 88 
 82 
 74 
 
 64 
 
 10 
 
 73 
 
 72 
 
 7i 
 
 70 
 
 69 
 
 68 
 
 67 66 
 
 65 
 
 ii 
 
 12 
 13 
 14 
 
 63 
 
 53 
 40 
 
 2 7 
 
 62 
 5i 
 39 
 26 
 
 61 
 
 5 
 38 
 24 
 
 60 
 49 
 37 
 23 
 
 59 
 48 
 
 35 
 
 22 
 
 58 
 47 
 34 
 20 
 
 57 
 45 
 33 
 19 
 
 56 
 
 44 
 3i 
 17 
 
 55 
 43 
 
 3 
 16 
 
 54 
 42 
 29 
 14 
 
 15 
 
 13 
 
 ii 
 
 10 
 
 08 
 
 06 
 
 5 
 
 3 
 
 02 
 
 00 
 
 99 
 
 16 
 17 
 18 
 19 
 
 0-99897 
 80 
 62 
 43 
 
 95 
 78 
 60 
 4i 
 
 94 
 77 
 59 
 39 
 
 92 
 
 75 
 57 
 37 
 
 9 
 73 
 
 55 
 35 
 
 89 
 7 1 
 53 
 33 
 
 87 
 
 70 
 
 5i 
 3i 
 
 85 
 
 68 
 
 49 
 29 
 
 84 
 66 
 
 47 
 
 27 
 
 82 
 64 
 45 
 25 
 
 20 
 
 2 3 
 
 21 
 
 i9 
 
 17 
 
 15 
 
 13 
 
 ii 
 
 08 
 
 06 04 
 
 21 
 22 
 
 23 
 24 
 
 02 
 0-99780 
 
 5 6 
 
 32 
 
 00 
 
 77 
 54 
 30 
 
 98 
 75 
 5 2 
 27 
 
 95 
 73 
 49 
 25 
 
 93 
 7i 
 
 47 
 
 22 
 
 9i 
 68 
 
 45 
 
 20 
 
 89 
 66 
 
 42 
 i7 
 
 86 
 64 
 40 
 15 
 
 84 
 61 
 
 37 
 
 12 
 
 82 
 59 
 35 
 10 
 
 2 5 
 
 07 
 
 5 
 
 02 
 
 99 
 
 97 
 
 94 
 
 92 89 86 
 
 8j 
 
 26 
 
 11 
 
 29 
 
 0-99681 
 
 11 
 
 o' 99597 
 
 78 
 5i 
 23 
 94 
 
 7 6 
 48 
 2O 
 91 
 
 7 l 
 46 
 
 17 
 
 88 
 
 70 
 
 43 
 15 
 85 
 
 67 
 40 
 
 12 
 
 82 
 
 65 
 37 
 09 
 
 79 
 
 62 59 
 
 34 32 
 06 03 
 
 7 6 73 
 
 57 
 29 
 
 00 
 
 70 
 
 30 
 
 67 
 
 64 
 
 61 
 
 58 
 
 55 
 
 S 2 
 
 49 
 
 46 
 
 43 
 
 40 
 
 31 
 32 
 
 33 
 34 
 
 37 
 05 
 o'99473 
 40 
 
 34 
 
 02 
 70 
 36 
 
 3 
 
 99 
 66 
 
 33 
 
 2 7 
 96 
 63 
 3 
 
 24 
 
 ? 2 
 60 
 
 26 
 
 21 
 
 89 
 56 
 23 
 
 18 
 86 
 53 
 
 20 
 
 15 
 83 
 
 5 
 16 
 
 12 
 
 79 
 46 
 
 13 
 
 08 
 76 
 
 43 
 09 
 
 35 
 
 06 
 
 02 
 
 99 
 
 95 
 
 92 
 
 89 
 
 85 
 
 82 
 
 78 
 
 75 
 
 *Watson. 
 
APPENDIX II. 
 
 385 
 
 TABLE VI. 
 
 VOLUME OF ONE GRAM OF WATER.* 
 
 (Thiesen, Scheel, and Diesselhorst, in Wiss. Abh. d. Phys. Tech. Reichsanstalt, Vol. 
 II, 1895; Vol. Ill, 1900; and Vol. IV, 1904.) 
 
 Tem- 
 pera- 
 ture. 
 
 
 
 i 
 
 2 
 
 3 
 
 4 
 
 5 
 
 6 
 
 7 
 007 
 
 8 
 
 9 
 
 
 
 I '00013 
 
 
 
 007 
 
 00 3 
 
 OOI 
 
 000 
 
 OOI 
 
 003 
 
 012 
 
 019 
 
 10 
 
 027 
 
 037 
 
 048 
 
 060 
 
 073 
 
 087 
 
 103 
 
 120 
 
 138 
 
 157 
 
 20 
 
 177 
 
 199 
 
 221 
 
 244 
 
 268 
 
 294 
 
 320 
 
 347 
 
 376 
 
 405 
 
 30 
 
 435 
 
 466 
 
 497 
 
 530 
 
 563 
 
 598 
 
 633 
 
 669 
 
 706 
 
 743 
 
 40 
 
 782 
 
 821 
 
 861 
 
 901 
 
 943 
 
 985 
 
 028 
 
 072 
 
 116 
 
 ^62 
 
 5 
 
 i '01207 
 
 254 
 
 301 
 
 349 
 
 398 
 
 448 
 
 498 
 
 548 
 
 600 
 
 652 
 
 60 
 
 705 
 
 758 
 
 812 
 
 867 
 
 923 
 
 979 
 
 036 
 
 093 
 
 iS 1 
 
 210 
 
 70 
 
 i '02270 
 
 330 
 
 390 
 
 45 2 
 
 5M 
 
 576 
 
 6 39 
 
 73 
 
 768 
 
 833 
 
 80 
 
 899 
 
 9 6 5 
 
 032 
 
 099 
 
 768 
 
 237 
 
 306 
 
 376 
 
 447 
 
 5^8 
 
 90 
 
 1-03590 
 
 663 
 
 736 
 
 810 
 
 884 
 
 959 
 
 035 
 
 in 
 
 188 
 
 265 
 
 100 11-04343 
 
 422 
 
 5oi 
 
 
 
 
 
 
 
 
 *i gram of ice has a volume of i .09008 cm. 3 
 
 TABLE VII. 
 
 DENSITY OF MERCURY.* 
 (At the Temperature (t) on the Normal Hydrogen Scale.) 
 
 t 
 
 Density 
 
 t 
 
 Density 
 
 t 
 
 Density 
 
 t 
 
 Density 
 
 t 
 
 Density 
 
 t 
 
 Density 
 
 
 
 I 3-595 
 
 
 
 
 
 
 
 
 
 
 
 I 
 
 5925 
 
 6 
 
 .5802 
 
 ii 
 
 5 6 79 
 
 16 
 
 555 6 
 
 21 
 
 5433 
 
 26 
 
 5310 
 
 2 
 
 .5901 
 
 7 
 
 5777 
 
 12 
 
 5 6 54 
 
 J 7 
 
 5531 
 
 22 
 
 .5408 
 
 27 
 
 .5286 
 
 3 
 
 .5876 
 
 8 
 
 5753 
 
 13 
 
 .5629 
 
 18 
 
 5507 
 
 2 3 
 
 5384 
 
 28 
 
 .5261 
 
 4 
 
 5851 
 
 9 
 
 5728 
 
 14 
 
 5 6 5 
 
 !9 
 
 5482 
 
 24 
 
 5359 
 
 29 
 
 5 2 37 
 
 5 
 
 13-5827 
 
 10 
 
 I3-5703 
 
 15 
 
 i3-558i 
 
 20 
 
 J3-5457 
 
 2 5 
 
 13-5335 
 
 3 
 
 13.5212 
 
 *Thiesen, Guillaume, International Congress of Physics, Paris, 1900. 
 25 
 
3 86 
 
 EXPERIMENTAL CHEMISTRY. 
 
 TABLE VIII. BAROMETRIC CORRECTIONS. 
 
 (a) REDUCTION or BAROMETER READINGS TO o. CORRECTIONS FOR THE EXPAN- 
 SION OF THE MERCURY AND THE SCALE. 
 The following corrections* are to be subtracted. 
 
 Tem- 
 pera- 
 ture. 
 
 BRASS SCALE. 
 
 GLASS SCALE. 
 
 OBSERVED HEIGHT. 
 
 OBSERVED HEIGHT. 
 
 c. 
 
 740 
 
 750 
 
 760 
 
 770 
 
 780 
 
 740 
 
 750 
 
 7 6o 
 
 770 
 
 7 80 
 
 I 
 
 2 
 3 
 
 4 
 
 mm. 
 
 12 
 
 mm. 
 
 . 12 
 25 
 
 37 
 49 
 
 mm. 
 
 . 12 
 
 25 
 37 
 SO 
 
 mm. 
 13 
 25 
 38 
 50 
 
 mm. 
 13 
 
 5i 
 
 mm. 
 
 38 
 Si 
 
 mm. 
 13 
 .26 
 39 
 52 
 
 mm. 
 13 
 .26 
 39 
 53 
 
 mm. 
 13 
 27 
 .40 
 S3 
 
 mm. 
 13 
 
 27 
 41 
 54 
 
 5 
 
 . 60 
 
 .61 
 
 .62 
 
 63 
 
 .64 
 
 .64 
 
 .65 .66 .67 
 
 .68 
 
 6 
 
 I 
 
 9 
 
 .72 
 8s 
 97 
 1-09 
 
 73 
 .86 
 .98 
 
 I- 10 
 
 74 
 87 
 99 
 
 I 12 
 
 75 
 .88 
 
 I-OI 
 
 I-I3 
 
 .76 
 .89 
 
 1-02 
 
 77 
 .90 
 
 1-02 
 
 .78 
 9i 
 1-04 
 i-i7 
 
 79 ; -80 
 92 ! -93 
 I- 05 ! 1-07 
 
 I. l8 1-20 
 
 81 
 
 95 
 i. 08 
 
 I 21 
 
 10 
 
 I- 21 
 
 I- 22 
 
 1-24 
 
 1.26 
 
 I- 27 
 
 1.28 
 
 1.30 1-31 1-33 
 
 i-35 
 
 ii 
 
 12 
 13 
 14 
 
 1-33 
 1-45 
 1-57 
 1-69 
 
 35 
 
 i'47 
 i'59 
 1-71 
 
 1-49 
 
 1-61 
 i-73 
 
 1.38 
 1^76 
 
 1-40 
 'S3 
 1-65 
 1-78 
 
 I-4I 
 
 i-53 
 1.66 
 1.79 
 
 1-43 
 1-56 
 1-69 
 1.81 
 
 i-45 
 1-58 
 1-71 
 
 1-84 
 
 1.46 
 i. 60 
 1-73 
 1.86 
 
 1.48 
 1.62 
 i-75 
 1.89 
 
 IS 1-81 
 
 1-83 
 
 1.86 
 
 1-88 
 
 1-91 
 
 1.92 
 
 1-94 
 
 1-97 2-OO 
 
 2 O2 
 
 O 00<I Ov 
 
 1-93 
 2-05 
 2- 17 
 2-29 
 
 1-96 
 2.08 
 
 2- 2O 
 2-32 
 
 1-98 
 
 2- 10 
 
 2.23 
 2-35 
 
 2-01 
 
 2.13 
 
 2- 26 
 2.38 
 
 2.03 
 2-16 
 
 2-29 
 2.41 
 
 2-05 
 2-17 
 2-30 
 2-43 
 
 2-07 
 
 2- 20 
 2-33 
 2-46 
 
 2-10 2-13 
 2-23 2.26 
 
 2.36 2-39 
 2-49 2.53 
 
 2-16 
 
 2- 29 
 
 ::0 
 
 2O 
 
 2.41 
 
 2-44 
 
 2-47 
 
 2-SI 
 
 2-54 
 
 2-56 
 
 2-59 2-62 i 2-66 2-69 
 
 21 
 22 
 23 
 
 24 
 
 2-53 
 2-65 
 2-77 
 2-89 
 
 2.56 
 2.69 
 
 2.81 
 2-93 
 
 2-60 
 2-72 
 2-84 
 2-97 
 
 2.63 
 2.76 
 
 2-88 
 3-oi 
 
 2.67 
 2-79 
 2-92 
 3 -05 
 
 2-68 
 2.81 
 2.94 
 3-06 
 
 2-72 
 2.85 
 2.98 
 3- ii 
 
 2- 76 
 
 2.89 
 
 3-02 
 
 3-15 
 
 2.79 
 2.92 
 3-o6 
 3-19 
 
 2.83 
 
 2.96 
 3-10 
 
 3-23 
 
 25 
 
 3-oi 
 
 3-05 
 
 3-09 
 
 3-13 
 
 3-17 
 
 3-19 
 
 3-23 
 
 3-28 
 
 3-32 
 
 3-36 
 
 to to to to 
 
 VO 00<I Ov 
 
 3- 13 
 3-25 
 3-37 
 3-49 
 
 3-17 
 3-29 
 3-41 
 3-54 
 
 3-21 
 3-34 
 3-46 
 3-58 
 
 3-26 
 3-38 
 3-Si 
 3-63 
 
 3-33 
 3-42 
 
 III 
 
 3-32 
 3-45 
 3-57 
 3-70 
 
 3-36 
 3-49 
 3-62 
 3-75 
 
 3-41 
 3-54 
 3-67 
 3-8o 
 
 3-45 
 3-59 
 3-72 
 3-85 
 
 3-50 
 3-63 
 3-77 
 3-90 
 
 30 
 
 3-6i 
 
 3-66 
 
 3-71 
 
 3-75 
 
 3-8o 
 
 3-83 
 
 3-88 
 
 3-93 
 
 3-98 
 
 4-03 
 
 31 
 32 
 
 33 
 
 34 
 
 3-73 
 3-85 
 3-97 
 4.09 
 
 3-78 
 3-90 
 4.02 
 4.14 
 
 3-83 
 3-95 
 4-07 
 4.20 
 
 3-88 
 4.00 
 4-13 
 4-25 
 
 3-93 
 4-05 
 4-18 
 4-31 
 
 3-95 
 4-08 
 
 4-21 
 
 4-33 
 
 4-01 
 4-14 
 4-26 
 4-39 
 
 4.06 
 4-19 
 4-32 
 4-45 
 
 4-n 
 4-25 
 4-38 
 4-51 
 
 4.17 
 4-30 
 4-43 
 4-57 
 
 35 
 
 4.21 
 
 4-26 
 
 4-32 
 
 4-38 
 
 4-43 
 
 4.46 
 
 4-52 
 
 4-58 
 
 4-65 
 
 4.71 
 
 *The coefficient of expansion of mercury is 0.0001813. If j3 is the coefficienct of expansion 
 of the scale, an amount equal to (o.oooi8i3-j3)ht must be subtracted from the observed height 
 (h) to obtain the reading if the scale and mercury were reduced to o C. (Glass, ft = .000009; 
 brass, fi = . 00002; steel, )S = . 00012.) 
 
 (b) CORRECTION FOR VARIATION IN g. 
 
 Latitude 
 
 730 
 
 740 
 
 750 
 
 760 
 
 770 
 
 780 
 
 790 
 
 
 mm. 
 
 mm. 
 
 mm. 
 
 mm. 
 
 mm. 
 
 mm. 
 
 mm. 
 
 45 
 
 
 
 o 
 
 
 
 o 
 
 o 
 
 
 
 
 
 40 or 50 
 35 or 55 
 
 32 
 65 
 
 III 
 
 
 :i? 
 
 35 
 .68 
 
 111 
 
 35 
 .70 
 
APPENDIX II. 
 
 387 
 
 TABLE IX. 
 
 REDUCTION TO VACUUM OF WEIGHINGS MADE IN AIR. 
 
 If a body of density (D) has an apparent weight of (G) grams when weighed in 
 in the air, its weight reduced to vacuum is G + Gk grams, k is computed for air of 
 density .0012 and for brass weights of density 8.4, and platinum weights, 21.5. 
 
 f .0012 _ .0012 \ 
 
 Go = G I ~D~ ^T~J" G(l+k) 
 
 
 Correction in 
 
 Milligrams. 
 
 Density of Body. D. 
 
 Brass Weights. 
 
 Platinum Weights. 
 d. = 21 '5. 
 
 
 k. 
 
 k. 
 
 0-50 
 
 + 2-26 
 
 + 2-34 
 
 '55 
 
 2 '04 
 
 13 
 
 60 
 
 86 
 
 '94 
 
 65 
 
 70 
 
 '79 
 
 70 
 
 '57 
 
 66 
 
 '75 
 
 46 
 
 '55 
 
 80 
 
 36 
 
 '44 
 
 '85 
 
 27 
 
 36 
 
 90 
 
 
 28 
 
 '95 
 
 ' 12 
 
 ' 21 
 
 '0 
 
 06 
 
 '14 
 
 ' i 
 
 '95 
 
 '04 
 
 '2 
 
 86 
 
 '94 
 
 '3 
 
 78 
 
 87 
 
 4 
 
 "7 1 
 
 80 
 
 '5 
 
 66 
 
 '75 
 
 6 
 
 61 
 
 69 
 
 '7 
 
 ;s6 
 
 65 
 
 8 
 
 
 62 
 
 1-9 
 
 '49 
 
 58 
 
 2 '0 
 
 46 
 
 '54 
 
 2'5 
 
 '34 
 
 '43 
 
 3'o 
 
 26 
 
 '34 
 
 3'5 
 
 '20 
 
 29 
 
 4'o 
 
 16 
 
 24 
 
 5 
 
 ' 10 
 
 19 
 
 6-0 
 
 06 
 
 '14 
 
 8-0 
 
 + '01 
 
 09 
 
 IO'O 
 
 '02 
 
 06 
 
 15' 
 
 "06 
 
 '03 
 
 20 "o 
 
 'OS 
 
 004 
 
EXPERIMENTAL CHEMISTRY. 
 
 TABLE X. 
 
 Volume of a Glass Vessel at Various Temperatures. 
 Correction Factors for Calibrating Glass Vessels. 
 
 In calibrating glass vessels by weighing them filled with water (or 
 mercury) at a known t, it is necessary to correct for the effect of tempera- 
 ture on the density of the liquid, the buoyant effect of the air, and the 
 effect of temperature on the volume of the flask. To ascertain the true 
 volume (in cm. 3 ) of a flask filled with distilled water at a temperature of 
 20 C., multiply its apparent weight (in grams) by 1.0028. These "fac- 
 tors" assume the use of brass weights in air of density .0012 (approx.), 
 and that the coefficient of cubical expansion of glass is .000025. 
 
 /. 
 
 VOLUME- WATER. 
 
 /. 
 
 VOLUME-MERCUR . 
 
 o 
 
 ccm. 
 
 o 
 
 ccm. 
 
 5 
 
 i.ooi 49 
 
 5 
 
 0.073 6 47 
 
 6 
 
 .001 49 
 
 6 
 
 .073 658 
 
 7 
 
 .001 50 
 
 7 
 
 .073 669 
 
 8 
 
 ooi 53 
 
 8 
 
 .073 681 
 
 9 
 
 .001 57 
 
 9 
 
 073 6 93 
 
 10 
 
 i.ooi 63 
 
 IO 
 
 0.073 704 
 
 ii 
 
 .001 70 
 
 ii 
 
 .073 716 
 
 12 
 
 .001 78 
 
 12 
 
 .073 727 
 
 13 
 
 .001 88 
 
 13 
 
 073 739 
 
 14 
 
 .001 99 
 
 14 
 
 073 750 
 
 15 
 
 I. O02 IO 
 
 15 
 
 0.073 762 
 
 16 
 
 .002 23 
 
 16 
 
 073 773 
 
 17 
 
 .002 38 
 
 17 
 
 073 785 
 
 18 
 
 .002 53 
 
 18 
 
 073 797 
 
 19 
 
 .002 70 
 
 19 
 
 .073 808 
 
 20 
 
 I .002 87 
 
 20 
 
 0.073 820 
 
 21 
 
 .003 06 
 
 21 
 
 073 831 
 
 22 
 
 .003 26 
 
 22 
 
 073 843 
 
 2 3 
 
 .003 47 
 
 2 3 
 
 073 854 
 
 24 
 
 .003 69 
 
 24 
 
 .073 866 
 
 25 
 
 1.003 9 2 
 
 25 
 
 0.073 877 
 
 26 
 
 .004 15 
 
 26 
 
 .073 889 
 
 27 
 
 .004 39 
 
 27 
 
 .073 900 
 
 28 
 
 .004 65 
 
 28 
 
 .073 912 
 
 2 9 
 
 .004 92 
 
 29 
 
 073 923 
 
 30 
 
 1.005 I 9 
 
 30 
 
 0.073 935 
 
APPENDIX II. 
 
 .589 
 
 TABLE XI. 
 
 VAPOR TENSION OF WATER. 
 Between o C and 35 C. 
 
 w 
 w 
 M 
 
 o 
 
 H 
 
 Q 
 
 VAPOR TENSION IN MILLIMETRES OF MERCURY. 
 
 TENTHS. 
 
 .0 
 
 ' 
 
 .2 
 
 3 
 
 4 
 
 5 
 
 -6 
 
 -7 
 
 .8 
 
 9 
 
 
 
 4-57 
 
 4.60 
 
 4.64 4-67 
 
 4- 70 
 
 4-74 
 
 4-77 
 
 4-80 
 
 4-84 4-87 
 
 i 
 
 2 
 
 3 
 
 4 
 
 4.91 
 5-27 
 5-66 
 6-07 
 
 4-94 
 5-31 
 5- 70 
 6. ii 
 
 4.98 
 5-35 
 5-74 
 6.15 
 
 5-02 
 
 5-39 
 5-78 
 
 6.20 
 
 5-05 
 5-42 
 5-82 
 6.24 
 
 5-09 
 5-46 
 5-86 
 6.28 
 
 5-12 
 
 5-50 
 5-94 
 6-33 
 
 5-i6 
 5-54 
 5-94 
 6-37 
 
 5-20 
 
 5-58 
 
 5-99 
 6-42 
 
 5-23 
 5-62 
 6-03 
 6.46 
 
 5 
 6 
 
 2 
 
 9 
 
 6-Si 6.55 
 
 6.60 
 
 6.64 
 
 6.69 
 
 6-74 
 
 6.78 
 
 6-83 
 
 6-88 6.92 
 
 6.97 
 7-47 
 7-99 
 8-55 
 
 7.02 
 7-52 
 8. os 
 8.61 
 
 7-07 
 7-57 
 8-io 
 8.66 
 
 7- 12 
 7.62 
 8.15 
 
 8-72 
 
 7- i? 
 7-67 
 
 8.2! 
 8.78 
 
 7.22 
 7-72 
 8-27 
 8.84 
 
 7-26 
 7-78 
 8-32 
 8.90 
 
 7-3i 
 7-83 
 8-38 
 8.96 
 
 7-36 7-42 
 7-88 7-94 
 8-43 8-49 
 9 02 9-08 
 
 IO 
 
 9*14 
 
 9- 20 
 
 9.26 9.32 
 
 9-39 
 
 9-45 
 
 9-51 
 
 9-58 
 
 9-64 9.70 
 
 ii 
 
 12 
 13 
 
 14 
 
 9-77 
 10-43 
 ii. 14 
 11.88 
 
 9-8 3 
 
 io- 50 
 
 II. 21 
 IL96 
 
 9.90 
 10-57 
 11.28 
 12.04 
 
 9.96 
 10-64 
 II-36 
 12-12 
 
 10-03 
 io- 71 
 H-43 
 12.19 
 
 10-09 
 10-78 
 11.50 
 12-27 
 
 io. 16 
 10-85 
 11.58 
 12-35 
 
 10.23 
 io 92 
 11.66 
 12-43 
 
 10-30 
 10-99 
 n-73 
 12-51 
 
 10-36 
 11-07 
 ii. 81 
 12-59 
 
 15 
 
 12-67 
 
 12- 76 
 
 12.84 
 
 12-92 
 
 13-00 
 
 13.09 
 
 13-17 
 
 13-25 
 
 13-34 
 
 13.42 
 
 16 
 
 ii 
 
 19 
 
 13-Si 
 14.40 
 15-33 
 16.32 
 
 13.60 
 14-49 
 15-43 
 16.42 
 
 13-68 
 14-58 
 15-52 
 16.52 
 
 13-77 ; 
 14-67 
 15-62 
 16.63 
 
 13-86 
 14-76 
 15-72 
 i6.73 
 
 13-95 
 14-86 
 15-82 
 16-83 
 
 14.04 
 14-95 
 15-92 
 16.94 
 
 14.12 
 
 14.21 
 
 14-30 
 15-23 
 
 l6- 22 
 17.26 
 
 16. 02 
 17-04 
 
 16. 12 
 17-15 
 
 20 
 
 I7-36 
 
 17-47 
 
 I7-58 
 
 17.69 
 
 17-80 
 
 17.91 
 
 18-02 
 
 18.13 
 
 18.24 
 
 18.35 
 
 21 
 22 
 23 
 
 24 
 
 18-47 
 19.63 
 20-86 
 22-15 
 
 18-58 
 19-75 
 20-98 
 
 22- 29 
 
 18-69 
 19-87 
 
 21. II 
 22-42 
 
 18.81 
 19-99 
 21.24 
 22-55 
 
 18-92 
 
 20- II 
 2L37 
 22.69 
 
 19-04 
 
 20- 24 
 21 -50 
 22.83 
 
 19. 16 
 20.36 
 21-63 
 22.96 
 
 19.27 
 20 48 
 
 21. 76 
 23.10 
 
 19-39 
 20.61 
 21-89 
 23-24 
 
 19-51 
 
 20-73 
 22-02 
 23-38 
 
 25 
 
 23-52 
 
 23-66 
 
 23-80 | 23.94 
 
 24.08 
 
 24-23 
 
 24-37 
 
 24-52 
 
 24-66 
 
 24-81 
 
 26 
 
 3 
 
 29 
 
 24.96 
 26-47 
 28-07 
 29.74 
 
 25. io 
 
 26.63 
 28.23 
 29.92 
 
 25-25 
 26-78 
 28.39 
 30.09 
 
 25-40 
 26-94 
 28-56 
 30.26 
 
 25-55 
 27- io 
 28.73 
 30-44 
 
 25-70 
 27. 26 
 28-89 
 30.62 
 
 25-86 
 27.42 
 29.06 
 30-79 
 
 26.01 
 27-58 
 29-23 
 30-97 
 
 26.16 
 27.74 
 29-40 
 3I-I5 
 
 26-32 
 27.90 
 29-57 
 31-33 
 
 30 
 
 3i-5i 
 
 31.69 
 
 3I-87 
 
 32.06 
 
 32-24 
 
 32-43 
 
 32.61 
 
 32.80 
 
 32.99 
 
 33-i8 
 
 3i 
 
 32 
 
 33 
 
 34 
 
 33-37 
 35-32 
 37-37 
 39-52 
 
 33-56 
 35-52 
 
 37-58 
 
 39-74 
 
 33-75 
 35-72 
 37-79 
 39-97 
 
 33-94 
 35-92 
 38-00 
 40- 19 
 
 34-14 
 36-13 
 38-22 
 40.41 
 
 34-33 
 36-33 
 38-43 
 40.64 
 
 34-53 
 36-54 
 38-65 
 40.86 
 
 34-72 
 36-74 
 38.87 
 41.09 
 
 34-92 
 36-95 
 39-o8 
 41-33 
 
 35-12 
 37-16 
 39-30 
 41-55 
 
 35 
 
 41-78 
 
 42 02 
 
 42-25 
 
 42-48 
 
 42-72 
 
 42.96 
 
 43- 19 
 
 43-43 
 
 43-6? 
 
 43-92 
 
390 
 
 EXPERIMENTAL CHEMISTRY. 
 
 TABLE XII. 
 VAPOR TENSION OF WATER. 
 
 FROM oC. TO iooC. 
 
 Temp 
 
 Tension 
 
 Temp 
 
 Tension 
 
 Temp 
 
 Tension 
 
 Temp 
 
 Tension 
 
 o 
 
 mm. 
 4-569 
 
 2 5 
 
 mm. 
 23-5!7 
 
 5 
 
 mm. 
 91.98 
 
 75 
 
 mm. 
 
 288.76 
 
 i 
 
 4.909 
 
 26 
 
 24.956 
 
 5i 
 
 96.66 
 
 76 
 
 301.09 
 
 2 
 
 5.272 
 
 27 
 
 26.471 
 
 52 
 
 IOI -55 
 
 77 
 
 313-85 
 
 3 
 
 5-658 
 
 28 
 
 28.065 
 
 53 
 
 106.65 
 
 78 
 
 327- 5 
 
 4 
 
 6.069 
 
 29 
 
 29.744 
 
 54 
 
 111.97 
 
 79 
 
 340-73 
 
 5 
 
 6.507 
 
 30 
 
 31- 5 1 
 
 55 
 
 117.52 
 
 80 
 
 354.87 
 
 6 
 
 6.972 
 
 3 1 
 
 33-37 
 
 56 
 
 123.29 
 
 81 
 
 369-51 
 
 7 
 
 7.466 
 
 32 
 
 35-32 
 
 57 
 
 129.31 
 
 82 
 
 384.64 
 
 8 
 
 7.991 
 
 33 
 
 37-37 
 
 58 
 
 I35-58 
 
 83 
 
 400.29 
 
 9 
 
 8.548 
 
 34 
 
 39-52 
 
 59 
 
 142. 10 
 
 84 
 
 416.47 
 
 10 
 
 9.140 
 
 35 
 
 41.78 
 
 60 
 
 148.88 
 
 85 
 
 433- 19 
 
 ii 
 
 9.767 
 
 36 
 
 44.16 
 
 61 
 
 155-95 
 
 86 
 
 450.47 
 
 12 
 
 10.432 
 
 37 
 
 46.65 
 
 62 
 
 163.29 
 
 87 
 
 468.32 
 
 13 
 
 11.137 
 
 38 
 
 49.26 
 
 63 
 
 170.02 
 
 88 
 
 486.76 
 
 14 
 
 11.884 
 
 39 
 
 52.00 
 
 64 
 
 178.86 
 
 89 
 
 505-81 
 
 15 
 
 12.674 
 
 40 
 
 54.87 
 
 65 
 
 187. 10 
 
 90 
 
 525-47 
 
 16 
 
 i3-5 10 
 
 4i 
 
 57-87 
 
 66 
 
 I95-67 
 
 9i 
 
 545-77 
 
 17 
 
 14-395 
 
 42 
 
 61.02 
 
 67 
 
 204.56 
 
 92 
 
 566.71 
 
 18 
 
 15-330 
 
 43 
 
 64-3 1 
 
 68 
 
 213.79 
 
 93 
 
 588.83 
 
 J 9 
 
 16.319 
 
 44 
 
 67.76 
 
 69 
 
 223.37 
 
 94 
 
 610.64 
 
 20 
 
 17-363 
 
 45 
 
 71.36 
 
 70 
 
 233-3I 
 
 95 
 
 633.66 
 
 21 
 
 18.466 
 
 46 
 
 75-13 
 
 7 1 
 
 243-62 
 
 96 
 
 657.40 
 
 22 
 
 19.630 
 
 47 
 
 79.07 
 
 72 
 
 254-3 
 
 97 
 
 681.88 
 
 23 
 
 20.858 
 
 48 
 
 83.19 
 
 73 
 
 265.38 
 
 98 
 
 707-13 
 
 24 
 
 22. 152 
 
 49 
 
 87.49 
 
 74 
 
 276.87 
 
 99 
 
 733.i6 
 
 *Taken from Ostwald's Manual of Physico-Chemical Measurements. 
 
APPENDIX II. 
 
 391 
 
 TABLE XIII. 
 
 VAPOR PRESSURE OF MERCURY. 
 (Ramsay and Young, Jour, of the Chem. Soc., 1886, p. 37.) 
 
 Temperature. 
 
 Vapor Pressure, 
 mm. 
 
 Temperature. 
 
 Vapor Pressure, 
 mm. 
 
 40 
 
 o'OoS 
 
 1 80 
 
 S'535 
 
 50 
 
 015 
 
 190 
 
 12-137 
 
 60 
 
 '029 
 
 200 
 
 17-015 
 
 70 
 
 052 
 
 210 
 
 23-482 
 
 80 
 
 092 
 
 220 
 
 31-957 
 
 90 
 
 160 
 
 230 
 
 42-919 
 
 100 
 
 270 
 
 240 
 
 5 6 "9!9 
 
 no 
 
 '445 
 
 250 
 
 74-592 
 
 120 
 
 719 
 
 260 
 
 96-661 
 
 130 
 
 i'i37 
 
 2 7 
 
 123-905 
 
 I4O 
 
 i"7 6 3 
 
 280 
 
 !57"37 8 
 
 ISO 
 
 2-684 
 
 290 
 
 198-982 
 
 1 60 
 
 4'oi3 
 
 300 
 
 246-704 
 
 170 
 
 5'94 
 
 
 
392 
 
 EXPERIMENTAL CHEMISTRY. 
 
 TABLE XIV. 
 
 BOILING TEMPERATURE (0 OF WATER AT BAROMETRIC PRESSURE (b). 
 (After Regnault.) 
 
 b 
 
 t 
 
 b 
 
 / 
 
 b 
 
 t 
 
 b 
 
 t 
 
 b 
 
 t 
 
 680 
 
 96.92 
 
 700 
 
 97.72 
 
 720 
 
 98.49 
 
 740 
 
 99.26 
 
 760 
 
 100.00 
 
 68 1 
 
 .96 
 
 01 
 
 75 
 
 21 
 
 53 
 
 41 
 
 .29 
 
 61 
 
 .04 
 
 682 
 
 97.00 
 
 02 
 
 79 
 
 22 
 
 57 
 
 42 
 
 33 
 
 62 
 
 .07 
 
 683 
 
 .04 
 
 3 
 
 .83 
 
 23 
 
 .61 
 
 43 
 
 37 
 
 63 
 
 . ii 
 
 684 
 
 .08 
 
 04 
 
 .87 
 
 24 
 
 65 
 
 44 
 
 .41 
 
 64 
 
 15 
 
 685 
 
 .12 
 
 5 
 
 .91 
 
 25 
 
 .69 
 
 45 
 
 44 
 
 65 
 
 .18 
 
 686 
 
 .16 
 
 06 
 
 95 
 
 26 
 
 .72 
 
 46 
 
 48 
 
 66 
 
 .22 
 
 687 .20 
 
 07 
 
 97-99 
 
 27 
 
 .76 
 
 47 
 
 5 2 
 
 67 
 
 .26 
 
 688 
 
 .24 
 
 08 
 
 98.03 
 
 28 
 
 .80 
 
 48 
 
 56 
 
 68 
 
 .29 
 
 689 
 
 .28 
 
 09 
 
 .07 
 
 29 
 
 .84 
 
 49 
 
 59 
 
 69 
 
 33 
 
 690 
 
 32 
 
 710 
 
 . ii 
 
 73 
 
 .88 
 
 75 
 
 63 
 
 77 
 
 .36 
 
 691 
 
 .36 
 
 ii 
 
 IS 
 
 3i 
 
 .92 
 
 5i 
 
 .67 
 
 7 1 
 
 .40 
 
 692 
 
 .40 
 
 12 
 
 .19 
 
 32 
 
 95 
 
 52 
 
 .70 
 
 72 
 
 44 
 
 693 
 
 44 
 
 13 
 
 .22 
 
 33 
 
 98.99 
 
 53 
 
 74 
 
 73 
 
 47 
 
 694 
 
 .48 
 
 14 
 
 .26 
 
 34 
 
 99 -3 
 
 54 
 
 78 
 
 74 
 
 5 1 
 
 695 
 
 52 
 
 15 
 
 30 
 
 35 
 
 .07 
 
 55 
 
 .82 
 
 75 
 
 55 
 
 696 
 
 -56 
 
 16 
 
 34 
 
 36 
 
 . ii 
 
 56 
 
 85 
 
 76 
 
 58 
 
 697 
 
 .60 
 
 17 
 
 38 
 
 37 
 
 .14 
 
 57 
 
 .89 
 
 77 
 
 .62 
 
 698 
 
 .64 
 
 18 
 
 .42 
 
 38 
 
 .18 
 
 58 
 
 93 
 
 78 
 
 65 
 
 699 
 
 .68 
 
 19 
 
 .46 
 
 39 
 
 .22 
 
 59 
 
 .96 
 
 79 
 
 .69 
 
 700 
 
 97.72 
 
 720 
 
 98.49 
 
 740 
 
 99.26 
 
 760 
 
 100 . OO 
 
 780 
 
 100.72 
 
 TABLE XV. 
 
 BOILING TEMPERATURE" (/) OF WATER AT A PRESSURE OF (a) ATMOSPHERES. 
 
 (Regnault.) 
 
 t 
 
 a 
 
 / 
 
 a 
 
 100 C. 
 
 I .000 
 
 180 C. 
 
 9.929 
 
 121 
 
 2 .025 
 
 189 
 
 12. 125 
 
 134 
 
 3.008 
 
 199 
 
 15.062 
 
 144 
 
 4.000 
 
 213 
 
 19.997 
 
 152 
 
 4.971 
 
 225 
 
 25-I25 
 
 159 
 
 5.966 
 
 239 
 
 27-534 
 
 I 7 I 
 
 8.036 
 
 
 
APPENDIX II. 
 
 393 
 
 4H 
 
 III 11 
 
 
 $ 
 
 'oS'o ^"o 
 
 >nCJ 
 
 
 If! 
 
 
 
 111 
 
 OMO o O H H H CM co 
 
 
 11? 
 
 JO O OOO -OOOOtoOO O O O O O 
 CO fO OOO O O ^t"CO t~> OO *O *O 'O^"O 
 
 
 Ill 
 
 S^o JOE* .g.coco^M^g . -^ .- . .OH ? 
 
 
 3 r3 . 
 
 1:5.1 
 
 to 
 
 O Tf " NO NO tOCO 
 
 H .M^MCOMCOCO.CO^^ ; ^ ; ^ ; J M 
 
 
 
 
 o o 
 
 
 3 CJ 4-J 
 
 
 
 1-S-s 
 
 oV S^NOv8No''d'^'^ ^ 
 
 
 3 a 
 
 ,5 ..^.^.Q ON 
 
 
 lit 
 
 o 
 
 ** O (N O O rO 
 tOtotoO HCNifO votoco NOIO to ^ONfOON 
 
 O O f^. t^^ ro fO f^ f^OGO O CO ON O ON fONO ON M CO ^f ^ O NO M CO 
 <N ON t^ t^OO ON <N CO OOO <N NO ON T|- NO CO ro M CO roJ5 rOCO W NO 
 
 u 
 
 
 
 o. 
 
 
 ^2 
 
 OJ 
 
 > 
 
 
 
 
 
 
 ti 
 
 
 : : gJU ::::::::::: :^| :::::: 
 
 (_ 
 
 s^ 
 
 
 SQ^IU ' TJ T3 .CL,W.. .. 
 
 * 
 
 
 
 *Critica 
 
394 
 
 EXPERIMENTAL CHEMISTRY. 
 
 TABLE XVII. 
 
 COMPOSITION OF THE AIR BY VOLUME. 
 (Average.) 
 
 Vols. per 1000. 
 
 Nitrogen 769 . 5000 
 
 Oxygen 206 . 5940 
 
 Aqueous vapor 14 . oooo 
 
 Argon * 
 
 Carbon dioxide 
 
 Hydrogen 
 
 Ammonia 
 
 Ozone 
 
 Nitric acid 
 
 9.3700 
 0.3360 
 o. 1900 
 o . 0080 
 0.0015 
 o . 0005 
 
 IOOO.OOOO 
 
 *The other four elements of the argon group constitute about .012 parts in a 1000. 
 Small quantities of solids are also present, e. g., ammonium nitrate, ammonium car- 
 bonate, sodium chloride, dust, etc. 
 
 TABLE XVIII. 
 
 DIFFUSION OF GASES. 
 (Graham, 1834 ) 
 
 
 Density 
 
 Square root 
 
 I 
 
 Actual 
 velocity of 
 
 Gas. 
 
 (air- 1) 
 
 of density 
 
 V Density 
 
 diffusion by 
 experiment 
 
 Hydrogen 
 
 o ' 06926 
 
 o ' 2632 
 
 3* 77Q4. 
 
 V8* 
 
 Marsh gas 
 
 O ' <"XQ 
 
 o ' 74.76 
 
 9-?7r 
 
 I ' 344. 
 
 Steam 
 
 o ' 62 3? 
 
 o ' 7806 
 
 '2664 
 
 
 Carbon monoxide 
 
 0*0678 
 
 0*0837 
 
 '016$ 
 
 I ' 1149 
 
 Nitrogen 
 
 O'Q7I3 
 
 w y">/ 
 o *o8c;6 
 
 '014.7 
 
 I '0143 
 
 Ethylene 
 
 O ' O?8 
 
 0*9889 
 
 ' OI 12 
 
 I ' 0191 
 
 Nitric oxide 
 
 O2Q 
 
 i * 1096 
 
 o ' 9808 
 
 
 Oxvgen . 
 
 ioc6 
 
 I ' O^I 1 
 
 O ' OCIO 
 
 0*0487 
 
 Sulphuretted hydrogen . 
 
 1912 
 
 I ' OQ 1 4. 
 
 o '9162 
 
 O'QC 
 
 Nitrous oxide 
 
 C27 
 
 I ' 23^7 
 
 0*8092 
 
 0*82 
 
 Carbon dioxide 
 
 * C2QOI 
 
 I '2^6i? 
 
 o ' 8087 
 
 0*812 
 
 Sulphur dioxide 
 
 2 '247 
 
 I ' 4QO I 
 
 0*667 I 
 
 o'68 
 
 
 
 
 
 
APPENDIX II. 
 
 395 
 
 TABLE XIX. 
 Specific Heat Atomic Heat. 
 
 Dulong-Petifs Law. A= 6 ' 4 ' 
 
 Elements. 
 
 S* 
 
 A 
 
 SxA 
 
 Lithium Li 
 
 O 04.1 
 
 7 
 
 6 6 
 
 Beryllium ... . .Be 
 
 o 408 
 
 
 7 7 
 
 Boron (amorphous) B . . . . 
 
 O.2C.4 
 
 II . 
 
 o / 
 
 2.8 
 
 Graphite \ 
 
 0,174 
 
 } 
 
 2.1 
 
 Diamond j 
 Sodium Na 
 
 0,143 
 O.27O 
 
 J 
 
 27 
 
 i-7 
 
 6.7 
 
 Magnesum. . Me . 
 
 O.24.CC 
 
 24.4 
 
 ">.Q 
 
 Aluminum Al 
 
 O,2O2 
 
 27 . 
 
 tr . cr 
 
 Silicon (cryst.) . . Si 
 
 o,i6c; 
 
 28 
 
 4.6 
 
 Phosphorus. . P 
 
 0,180 
 
 7T . 
 
 ^0 
 
 Sulphur (rhombic) S 
 
 0,178 
 
 72 . 
 
 c. . 7 
 
 Potassium K 
 
 o 166 
 
 TO 
 
 6.< 
 
 Calcium Ca 
 
 0,170 
 
 4O. 
 
 " 
 
 6.8 
 
 Chromium . . Cr 
 
 0,100 
 
 C2 . 
 
 C. . 2 
 
 Manganese Mn 
 
 O 122 
 
 err 
 
 6.7 
 
 Iron Fe 
 
 O,II2 
 
 ^.Q 
 
 6.7. 
 
 Cobalt Co 
 
 O IO7 
 
 CO 
 
 6.7. 
 
 Nickel . Ni 
 
 O 1 08 
 
 c8 7 
 
 6 4 
 
 Copper Cu . . 
 
 O,OQ7. 
 
 o u / 
 67.6 
 
 e. .0 
 
 Zinc Zn 
 
 O.OQ3 
 
 6? 4. 
 
 6 i 
 
 Gallium Ga 
 
 O O70 
 
 7O 
 
 r r 
 
 Germanium Ge 
 
 O,OC,7 
 
 72. C 
 
 cr .4 
 
 Arsenic (cryst.) As 
 
 O,o82 
 
 7cr . 
 
 6.2 
 
 Selenium (cryst ) Se 
 
 o 080 
 
 7O 2 
 
 6 4 
 
 Bromine (solid) . Br . .... 
 
 0,084 
 
 70.7 
 
 6 7 
 
 Zirconium Zr 
 
 0,066 
 
 QO. 6 
 
 6 o 
 
 Molybdenum Mo 
 
 0,072 
 
 06. 
 
 6.9 
 
 Ruthenium Ru 
 
 0,061 
 
 101 . 7 
 
 6.7. 
 
 Rhodium . . Rh 
 
 o.o<c8 
 
 103 . 
 
 6.0 
 
 Palladium Pd 
 
 o,oc,o 
 
 106. c 
 
 6.3 
 
 Silver . . . Ag .... 
 
 o,o<c6 
 
 107 .0 
 
 6.0 
 
 Cadmium Cd 
 
 O.OC.4 
 
 112 .4 
 
 6.0 
 
 Indium In 
 
 O OC.7 
 
 lie 
 
 6.4 
 
 Tin . Sn 
 
 0,0^4 
 
 IIQ. 
 
 6.4 
 
 Antimony. . . . Sb 
 
 O.OC.2 
 
 1 2O. 2 
 
 6.2 
 
 Tellurium Te 
 
 0.047 
 
 127.6 
 
 6.0 
 
 Iodine I 
 
 O OC.4 
 
 126.0 
 
 6.8 
 
 Lanthanum ... La 
 
 O,O4C. 
 
 138.0 
 
 6.2 
 
 
 
 
 
 * Richter. 
 
396 
 
 EXPERIMENTAL CHEMISTRY. 
 
 TABLE XIX. Continued. 
 
 Elements. 
 
 S* 
 
 A 
 
 SxA 
 
 Cerium Ce 
 
 0,04^ 
 
 I4O. 2 
 
 6.2 
 
 Didymium'j' Di 
 
 O.O4C 
 
 142 
 
 6 4 
 
 Tungsten \V 
 
 o 011 
 
 184 
 
 6 i 
 
 Osmium Os 
 
 ^J^OO 
 
 O,O3I 
 
 101 . 
 
 6.1 
 
 Iridium Ir 
 
 O.O32 
 
 103 
 
 6 3 
 
 Platinum Pt 
 
 O O32 
 
 IQ4 8 
 
 6 3 
 
 Gold ... Au 
 
 O,O^2 
 
 107 . 2 
 
 " o 
 
 6.3 
 
 Mercury (solid) Hg 
 
 O O32 
 
 2OO 
 
 6 4 
 
 Thallium Tl 
 
 O O3 3 
 
 2Q4 I 
 
 6 7 
 
 Lead . Pb 
 
 w >^oo 
 O,O3I 
 
 206 o 
 
 6 4 
 
 Bismuth Bi . 
 
 O O3O 
 
 208 
 
 6 3 
 
 Thorium . Th 
 
 O,O27 
 
 232 . <C 
 
 w o 
 
 6.4 
 
 Uranium Ur . . 
 
 o 027 
 
 238 q 
 
 6 4 
 
 
 
 
 
 * Richter. 
 
 fDidymium = Pr. and Nd. 
 
APPENDIX II. 
 
 397 
 
 TABLE XX. 
 Heat of Formation.* 
 
 (Arranged from Thomsen's Thermo-chemische Untersuchungen.) 
 (a) Chlorides, Bromides and Iodides. 
 
 Formula. 
 
 Chlorides. 
 
 Bromides. 
 
 Iodides. 
 
 (anhyd)f (aq.)J 
 
 (anhyd.) (aq.) 
 
 (anhyd.) 
 
 (aq.) 
 
 A1R 3 
 SbR 3 
 AsR s 
 BaR 3 
 BiR 3 
 CaR 2 
 CdR 2 
 CoR 2 
 CuR 
 CuR 2 
 AuR 
 AuR 3 
 HR 
 FeR 2 
 FeR 3 
 PbR 2 
 LiR 
 MgR 2 
 MnR 2 
 Hg 2 R 2 
 HgR 2 
 NiR 2 
 KR 
 AgR 
 NaR 
 SrR 2 
 TIR 
 TIR 3 
 SnR 2 
 SnR 4 
 ZnR 2 
 
 160900 
 91400 
 71400 
 194200 
 90600 
 183800 
 93200 
 76400 
 32800 
 51600 
 5800 
 22800 
 
 22OOO 
 82OOO 
 96OOO 
 82700 
 93800 
 I5IOOO 
 III9OO 
 652OO 
 
 54500 
 74OOO 
 105600 
 29300 
 97600 
 184500 
 48500 
 
 237700 
 
 129600 
 
 204900 
 
 70300 
 
 159300 
 
 
 47100 
 169400 
 
 
 12600 
 
 144000 
 
 196300 
 
 174400 
 
 
 2OI2OO 
 962OO 
 948OO 
 
 154900 
 75200 
 
 179400 
 75600 
 72900 
 
 107600 
 48900 
 
 14900 
 47900 
 42500 
 
 24900 
 32500 
 
 16200 
 
 62700 
 
 40800 
 
 
 
 
 27200 
 39300 
 999OO 
 127700 
 75900 
 IOI3OO 
 186900 
 I28OOO 
 
 8800 
 8400 
 
 
 55 
 
 
 28300 
 78000 
 
 6300 
 
 13100 
 47600 
 
 
 
 64400 
 79900 
 
 544oo 
 91300 
 165000 
 106100 
 
 39600 
 61200 
 
 
 76100 
 134600 
 75700 
 
 
 
 50900 
 4190 
 
 31100 
 25600 
 
 51200 
 93700 
 IOI200 
 
 
 
 71800 
 90200 
 
 41400 
 75000 
 
 953oo 
 27700 
 85700 
 157700 
 41400 
 
 80100 
 13800 
 69100 
 
 96500 
 195700 
 38400 
 89OOO 
 SlIOO 
 I57OOO 
 II280O 
 
 85600 
 173800 
 
 70300 
 143400 
 
 30100 
 
 56400 
 
 10500 
 
 80800 
 I272OO 
 97200 
 
 
 
 
 
 
 
 759oo 
 
 90900 
 
 49200 
 
 60500 
 
 * Smoothed values have been taken. 
 
 f Compound in usual state of aggregation. 
 
 | Compound formed in dilute aqueous solution. 
 
398 
 
 EXPERIMENTAL CHEMISTRY. 
 
 (b) Sulphates, Nitrates and Carbonates. 
 
 Formula. 
 
 Sulphates. 
 
 Formula. 
 
 Nitrates. 
 
 (anhyd.)| (aq.) 
 
 (anhyd.) | (aq.) 
 
 A1 2 (S0 4 ) 3 
 BaS0 4 
 CaSO 4 
 CdS0 4 
 CoS0 4 
 CuSO 4 
 H 2 S0 4 
 
 
 878900 
 
 
 
 
 337500 
 33240 
 
 22IIOO 
 233900 
 182500 
 192900 
 24OOOO 
 2l62OO 
 334100 
 3O22OO 
 249800 
 
 Ba(N0 3 ) 2 
 Ca(NO 3 ) 2 
 Cd(N0 3 ) 2 
 Co(N0 3 ) 2 
 Cu(N0 3 ) 2 
 HNOo 
 Fe7Fo 3 ) 2 
 Pb(N0 3 ) 2 
 LiN0 3 
 Mg(N0 3 ) 2 
 Mn(NO 3 ) 2 
 HgN0 3 
 Hg(N0 3 ) 2 
 Ni(N0 3 ) 2 
 KN0 3 
 AgN0 3 
 NaNO 3 
 Sr(N0 3 ) 2 
 T1NO 3 
 Zn(N0 3 ) 2 
 
 219800 
 216700 
 
 I2I2OO 
 II9300 
 92900 
 41600 
 
 2I52OO 
 2207OO 
 115800 
 II4300 
 82200 
 49IOO 
 II9500 
 97900 
 II22OO 
 206300 
 147500 
 38900 
 67100 
 II32OO 
 III400 
 23300 
 IO62OO 
 2I52OO 
 48lOO 
 132300 
 
 336800 
 231800 
 230400 
 198300 
 2IO7OO 
 235600 
 
 FeSO 4 .7H 2 O 
 PbSO 4 
 Li 2 S0 4 
 MgS0 4 
 MnSO 4 
 
 NiS0 4 
 K 2 SO 4 
 Ag 2 S0 4 
 Na 2 SO 4 
 SrSO 4 
 T1 2 S0 4 
 ZnSO 4 
 
 105500 
 III 6OO 
 2I057O 
 
 340100 
 323000 
 263500 
 
 2293OO 
 337200 
 162700 
 32900 
 
 
 
 120760 
 II94OO 
 28700 
 III2OO 
 219800 
 58lOO 
 138200 
 
 344500 
 167200 
 328500 
 330800 
 22IOOO 
 23OOOO 
 
 2I27OO 
 248OOO 
 
 BaCO 3 
 CdCO 3 
 MnCO 3 
 Ag 2 C0 3 
 SrC0 3 
 
 28l3OO 
 l8l900 
 2IO80O 
 I2I300 
 279600 
 
 
 CaCO 3 
 PbC0 3 
 K 2 C0 3 
 Na 2 CO 3 
 
 269200 
 I682OO 
 279500 
 27IOOO 
 
 
 
 
 
 286OOO 
 276500 
 
 
 
 
 
 
 
APPENDIX II. 
 
 399 
 
 (c) Oxides, Hydroxides and Sulphides. 
 
 Formula. 
 
 Oxides. 
 
 Formula. 
 
 Hydroxides. 
 
 (anhyd.) 
 
 (aq.) 
 
 (anhyd.) (aq.) 
 
 As 2 O 3 
 BaO 
 
 CaO 
 CuO 
 Cu 2 O 
 H 2 O 
 (vapor) 
 (liquid) 
 
 20 C. 
 
 PbO 
 Li 2 
 HgO 
 Hg 2 
 K 2 
 Ag 2 
 Na 2 
 SrO 
 
 154600 
 130400 
 14500 
 37100 
 40800 
 
 57061 
 68360 
 
 50300 
 
 147000 
 158200 
 163300 
 
 Al(OH), 
 Ba(OH) 2 
 Cd,0,H 2 
 Ca,0 2 ,H 2 
 Co,O,H 2 O 
 Au 2 3 , 3 H 2 
 Fe,0,H 2 
 Fe 27 3 , 3 H 2 
 
 LLO,H 
 Mg,0,H 2 
 Mn,O,H 2 O 
 Ni,O,H 2 O 
 K,0,H 
 NaOH 
 Sr,0 2 ,H 2 
 Zn,O,H 2 O 
 
 296900 
 216300 
 65700 
 
 
 2266OO 
 
 231700 
 
 
 63400 
 13200 
 68200 
 191100 
 
 cal. 
 cal. 
 
 
 
 
 II74OO 
 
 166500 
 
 148900 
 94700 
 60800 
 103900 
 
 IO2OOO 
 216400 
 82600 
 
 22OOO 
 24860 
 99100 
 5900 
 99800 
 128400 
 
 
 
 
 164500. 
 
 116400 
 I I l8oO 
 226lOO 
 
 155200 
 157700 
 
 
 A1 2 S 3 
 BaS 
 CaS 
 CuS 
 Cu 2 S 
 H 2 S 
 PbS 
 Li 2 S 
 MgS 
 HgS 
 K 2 S 
 A g2 S 
 Na 2 S 
 SrS 
 
 124400 
 99500 
 92OOO 
 IOOOO 
 
 18200 
 2700 
 18400 
 
 
 BaS 2 H 2 
 CaS 2 H 2 
 CdS,nH 2 
 FeS,nH 2 O 
 Co,S,nH 2 O 
 Ni,S,nH 2 
 LiSH 
 MgS 2 H 2 
 ZnS,nH 2 O 
 KSH 
 
 NaSH 
 SrS 2 H 2 
 
 
 I24IOO 
 125300 
 
 107100 
 
 II02OO 
 
 
 32400 
 2l8oO 
 19700 
 17400 
 
 
 
 7300 
 
 
 
 II32OO 
 I IOOOO? 
 
 64100 
 IIO8OO 
 
 79600 
 
 16800 
 
 102400 
 
 33000 
 
 88400 ? 
 99200? 
 
 
 39600 
 
 111300 
 
 63100 
 
 58500 
 II9700 
 
 55000 
 
 101900 
 104700 
 
 
400 
 
 EXPERIMENTAL CHEMISTRY. 
 
 ! 
 
 .J 
 
 T3 .S 
 
 1 
 
 H I 
 
 sapufdjns 
 
 sapixQ 
 
 sapipoj 
 
 sapjxcupAjj 
 
 sapuonjjj 
 
 sapiuiojg 
 
 N O 
 
 "o *o 
 
 g 
 
 1 3 
 
 *3 "o 
 
 w to 
 
 cfl rt 
 
 C C 
 
 I! 1 
 
 K W -s 
 
 . .- o 
 
 <s<s s? 
 
APPENDIX II. 
 
 4OI 
 
 X) 
 
 ON 10 
 
 t 
 
 oo 
 
 in 
 
 00 *c!i 
 
 o o 
 
 to 
 
 t* O 
 
 
 ST 
 
 NO^ 
 
 CO? 
 
 N CO 
 
 80 
 
 H ^ 
 
 O 
 
 S- "* 
 
 8 
 
 
 N vo 
 00 
 
 VO 
 
 ss 
 
 
 
 
 HI 
 
 
 
 
 
 o o 
 
 Z" 
 
 co 
 
 H O 
 
 
 
 
 
 
 
 
 
 
 
 o o 
 
 rt 
 
 ON <N 
 
 w oo 
 
 ON 
 
 S"S 
 
 
 
 oo t^ 
 
 ON co 
 
 SoO 
 
 wo 
 
 IO *S) 
 
 00 
 
 M 
 
 
 NO^ * 
 
 O 
 
 ^ 
 
 o o 
 
 
 ff* 
 
 00 ON 
 t^ 
 
 Tf 
 
 ON NO 
 
 H 
 t 
 
 
 
 *^- ^f 
 
 co 
 
 00 
 
 00 M 
 10 
 
 o o 
 
 o o 
 
 CO co 
 vo 
 
 
 
 
 O 
 
 be 
 
 00 H 
 
 MNO 
 
 M H 
 
 -? 
 80 
 
 M 
 
 
 NO 
 
 J^NQ 
 
 
 
 
 CO 
 
 TJ-OO 
 
 O co 
 
 co O 
 
 
 M 
 
 s 
 
 vo 
 
 2* 
 
 00 T 
 TT 
 
 
 
 ^ ^t 
 
 t^. 
 
 NO T 
 <N 
 
 rr 
 
 NO 
 
 
 
 VO N 
 
 CO 
 
 CO t 
 
 o o 
 
 O 
 
 cd 
 
 ON 
 
 H TJ- 
 
 co <s 
 
 00 
 
 NO 
 
 
 
 
 oo w 
 
 
 J 
 
 vo O 
 
 t>. N 
 
 vo 
 
 O HI 
 
 CO 
 
 NO CO 
 VO Tf 
 
 "> CO 
 
 oS 
 
 
 t^ 
 
 H 
 
 
 
 
 H 
 
 p. 
 
 00 
 
 
 
 
 
 
 
 
 
 
 O 
 
 CO 
 
 
 
 w co 
 
 vo rf 
 
 CO 
 
 ON 
 
 N ON 
 
 NO 
 
 S8 
 
 
 
 l^ 
 
 S" 
 
 ON NO 
 t * 
 
 ON 
 
 
 
 CO 
 
 t^ 
 
 "? 
 
 
 
 CO 
 r-^0 
 t- 
 
 
 
 M *?, 
 
 
 
 
 
 
 ^ 
 
 85 
 
 
 
 o^: 
 
 o o 
 
 
 
 
 
 
 
 NO ^ 
 
 
 
 
 H 
 
 
 co O 
 
 00,0 
 
 33 
 
 ro H 
 
 oj 
 
 IN t~- 
 
 NO ON 
 
 TOO 
 
 M 6' 
 
 r^ <*5 
 
 Tf M 
 
 00 
 
 o o 
 
 r~ N 
 
 o o 
 
 
 
 o o 
 
 
 
 
 to 
 
 O 
 
 
 
 
 
 00 
 
 CO 
 
 
 
 O 
 
 CO 
 
 
 
 
 
 O 
 
 
 
 
 fl 
 
 vn 
 
 *> J7 
 
 
 
 
 Cv 
 
 ON^ 
 
 
 
 ^0 H 
 
 
 
 
 [H 
 
 CO 
 
 
 
 
 
 
 
 ON CO 
 
 NO H 
 
 CO 
 
 
 
 !-> 
 
 r- 
 
 
 
 T 
 
 ON HI 
 
 
 o o 
 
 o o 
 
 
 
 oi ro 
 t-^. 
 
 00 O 
 
 co O 
 
 
 
 
 
 H 
 
 TT 
 
 
 
 M O 
 
 T 
 
 
 NO o 
 
 H M 
 
 ^ 
 
 10 
 
 o= 
 
 
 
 VO 
 
 vo 
 
 0*<N 
 
 tj 
 
 9 - 
 
 ? 9 
 
 
 
 :? 
 
 8 o 
 
 VO 
 
 g' 
 
 M 
 
 < 
 
 00 
 
 
 
 
 
 
 ON H 
 
 CO 00 
 
 H 
 
 M 
 
 
 
 
 
 
 
 
 
 
 
 
 
 O 
 
 
 ON 
 
 
 
 r~- H 
 
 CO 
 
 
 
 
 CO f 
 
 T 
 
 rr 
 
 
 N ON 
 
 t^ 
 
 
 t-- CO 
 
 t- M 
 
 00 <N 
 %0 M 
 
 IH 00 
 NO 
 
 
 
 r- 
 
 CO 
 
 r 
 
 oo 
 
 5 
 
 CO 
 
 H 
 
 
 
 
 NO <s 
 
 NO 
 
 
 5 
 
 t^ 
 
 NO 
 
 r- 
 
 CO 
 
 
 CO vo 
 
 HI O 
 
 t * 
 
 ^00 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 CO 
 
 00 
 
 H 
 
 
 00 
 
 ON 
 
 CO 
 
 
 M N 
 H 
 
 M 
 
 NO 
 
 
 M 
 
 M 
 
 ON ON 
 
 NO 
 
 00 NO 
 
 in 
 
 O 
 10 ^t 
 
 * 
 
 CJ 
 
 0000 
 
 c.,0 
 
 
 HI NO 
 
 H r^. 
 
 C-1 NO 
 
 ON 
 
 
 R" 
 
 vo * 
 
 t^O 
 
 
 Pi 
 
 NO 
 
 NO 
 
 
 ^r H 
 
 
 NO 
 
 CO 
 
 2 1 " 
 
 
 
 
 (_| 
 
 n 
 
 I I 
 
 fa 
 
 6 
 & 
 
 2 
 
 O 
 
 o 
 pq 
 
 o 
 
 a 
 
 
 
 
 
 <5 
 
 o 
 
 q 
 
 o 
 
 8 
 
402 
 
 EXPERIMENTAL CHEMISTRY. 
 
 TABLE XXIII. 
 Composition of Some of the Important Alloys. 
 
 Gold coin (U.S., Fr., Ger.) 
 
 Gold coin (Great Britain) 
 
 / Gold 900 
 
 \ Copper I0 o 
 
 / Gold 916.6 
 
 \? PPer 83.4 
 
 Silver coin (U. S.) 
 
 Silver coin (Great Britain) 
 
 Silverware (sterling) 
 
 Bronze 
 
 Gun-metal 
 
 Bell-metal 
 
 Speculum-metal 
 
 Aluminum bronze.. 
 
 Managanese bronze. . . 
 
 Phosphorus bronze ................. 
 
 Silicon bronze 
 
 Red brass. . 
 White brass 
 
 German silver 
 
 Silver 900 
 
 Copper ioo 
 
 Silver 925 
 
 Copper 75 
 
 Silver 925 
 
 Copper 75 
 
 Copper 93-5-95 
 
 Tin 6-4 
 
 Zinc 0.5-1 
 
 Copper ioo 
 
 Tin I0 
 
 Copper 78 
 
 Tin 22 
 
 67 
 
 r 33 
 
 topper 90-95 
 
 Aluminum 10-5 
 
 Copper 90 
 
 Managanese 10 
 
 Copper 90 
 
 Tin 9 
 
 Phosphorus 5~-8 
 
 Copper 90 
 
 Tin 
 
 Silicon 
 
 Copper 
 
 Zinc 
 
 Copper 
 
 Zinc 
 
 Copper 
 
 Zinc 
 
 Nickel 
 
 Tin.. 
 
 Britannia-metal. . 
 
 Hard pewter 
 
 9 
 
 8 
 
 90 
 
 10 
 
 65 
 
 35 
 
 5 
 
 25 
 
 25 
 
 ioo 
 
 Antimony 8 
 
 Bismuth i 
 
 Copper 4 
 
 Tin 92 
 
 Lead.. 8 
 
APPENDIX II. 
 
 TABLE XXIII. Continued. 
 
 403 
 
 Tin, 
 
 Solder (soft).. 
 Queen's metal, 
 
 Type-metal 
 
 so* pewter (Lead.:::: 
 
 Tin 
 
 Lead 
 
 Tin 
 
 Antimony. 
 Bismuth . 
 
 Lead 
 
 Tin 
 
 Lead 
 
 Antimony. 
 
 Fusible metal (" Rose's Metal"), .... l~ m '' ' ' 
 
 fuses at 94 C lu' ' 
 
 I Bismuth. . 
 
 (Tin 
 
 Fusible metal (" Wood's Alloy"), .... 
 fuses at 65 C 
 
 Babbitt metal,. 
 
 82 
 
 18 
 
 (varies) 50 
 
 (varies) 50 
 
 Lead 
 
 Bismuth. . 
 Cadmium. 
 
 Tin 
 
 Lead 
 
 Antimony. 
 Copper. . . 
 
 25 
 
 50 
 
 12 
 
 50 
 13 
 
 40 
 
 13 
 
 I 
 
404 EXPERIMENTAL CHEMISTRY. 
 
 TABLE XXIV. 
 Scale of Hardness. 
 
 Mineralogists have found it convenient to select a number of minerals 
 for the comparison of hardness. They designate the hardness on a 
 scale of 10. 
 
 No. i. Talc. Scratched easily by the finger-nail. 
 
 No. 2. Gypsum. Scratched with difficulty by the nail. 
 
 No. 3. Calcite. About as hard as a copper coin. 
 
 No. 4. Fluorite. Slightly harder than a copper coin. 
 
 No. 5. Apatite. Scratched easily by knife-point. 
 
 No. 6. Orthoclase. Scratched with difficulty by knife-point. 
 
 No. 7. Quartz. Is not scratched by knife-point. 
 
 No. 8. Topaz or beryl. 
 
 No. 9. Corundum. Similar to "emery." 
 
 No. 10. Diamond. 
 
 TABLE XXV. 
 Color Scale of Temperature. 
 
 This table is the result of an effort to interpret in terms of thermometric 
 readings, the common expressions used in chemistry in describing tem- 
 peratures. It is obvious that the values are only approximations. 
 
 Color 
 
 Temperature. 
 
 Incipient red heat. 
 
 5 oo-55o 
 
 Dark red heat. 
 
 6 5 o-75o 
 
 Bright red heat. 85o-95o 
 
 Yellowish-red heat. 
 
 1050-! 150 
 
 Incipient white heat. i25o-i35o 
 
 White heat. I45o-i55o 
 
APPENDIX II. 405 
 
 TABLE XXVI. 
 Indicators. 
 
 Phenolphthalein Alcohol solution i part solid in 100 
 
 (A weak acid.)* parts of 60 per cent, alcohol. Col- 
 
 C2oH 14 O 4 ? H- + C 20 H 13 O' 4 orless by acids; also by CO 2 ; red 
 
 (Colorless) (red) by alkalies. 
 
 Methylorange Water solution i part solid to 1000 
 
 (A weak base.) parts of water. Red by mineral 
 
 /N = N C 6 H 4 N(CH 3 ) 2 acids; not affected by CO 2 ; yellow 
 C 6 H 4 
 
 \SO 3 Na color by alkalies. 
 
 Litmus Water solution. Red by acids; blue by .alkalies. 
 
 Cochineal Alcoholic solution, 3 grm. of powder 
 
 in 400 parts H 2 O and 100 parts al- 
 cohol. Yellowish-red by acids; 
 violet by alkalies. 
 
 *Stieglitz, Jour. Amer. Chem. Soc. (Nov. 1903). 
 
 TABLE XXVII. 
 
 Soap- Bubble Solution No. I. 
 
 Newth suggests the following formula for preparing soap solutions: 
 "Ten grams of sodium oleate and 400 cm. 3 of distilled water are 
 placed in a stoppered bottle, and allowed to stand until the oleate has 
 dissolved, without warming; 100 cm. 3 of pure glycerin are then 
 added, and the mixture after being well shaken is allowed to stand 
 in the dark for a few days. The clear solution is then carefully 
 decanted or siphoned into a clean, stoppered bottle, one drop of 
 strong ammonia solution added. If kept in the dark and not ex- 
 posed to the air, this solution may be preserved for years." 
 
406 EXPERIMENTAL CHEMISTRY. 
 
 Soap-Bubble Solution No. II. 
 
 The following formula will be found to give excellent results: 
 Place about 100 grm. of shavings of the best castile soap in a liter 
 flask; add nearly a liter of distilled water. Shake the mixture until 
 a saturated solution is obtained, then allow it to settle. To two 
 volumes of the clear soap solution add one volume of glyerin. 
 Keep in a stoppered bottle in the dark. 
 
 TABLE XXVIII. 
 Graduated Solutions. 
 
 (a) Fehling's Solution. 
 
 This solution is used frequently in making a quantitative determina- 
 tion of sugar. As it decomposes on keeping, it is best kept (" pre- 
 pared") in the form of two separate solutions: (i) 34.639 grm. 
 of crystallized copper sulphate, made up to 500 cm. 3 with water; 
 (2) 173 grm. of Rochelle salt and 60 grm. of sodium hydroxide made 
 up to 500 cm. 3 with water. For use, equal volumes of the two solu- 
 tions are mixed. 10 cm. 3 of the deep blue solution thus obtained are 
 completely reduced by 0.05 grm. of dextrose, or by 0.0475 g rm - f 
 sucrose, after inversion. The color is discharged during the reduc- 
 tion. 
 
 (b) Nessler's Solution. 
 
 This solution is used in estimating ammonia (free). "Dissolve 17 
 grm. of mercuric chloride (pulverized) in 300 cm. 3 of water, and 35 
 grm. of potassium iodide in 100 cm. 3 of water. Pour the mercuric 
 chloride solution into the potassium iodide until a permanent red 
 precipitate is formed. Add a 20 per cent, solution of sodium hydrox- 
 ide until the volume of the mixed solution amounts to one liter. 
 Keep this solution in a cool, dark place portions of the solution 
 to be removed as needed. It is necessary to " sensitize " the solution 
 before it is ready for use; this is accomplished by adding slowly a 
 saturated solution of mercuric chloride, with constant stirring, 
 until a permanent red precipitate forms : Allow the solution to stand 
 until the solids subside, or filter. It is now ready for use and should 
 have a light, straw-yellow color. The solution deteriorates by 
 standing." 
 
APPENDIX II. 
 
 407 
 
 TABLE XXIX. 
 
 Percentage and Specific Gravity of Solutions at 15 C. 
 (a) Sulphuric Acid (Lunge and Isler). 
 
 Specific 
 Gravity. 
 
 Per cent. 
 H 2 SO 4 . 
 
 Specific 
 Gravity. 
 
 Per cent. 
 H 2 S0 4 . 
 
 Specific 
 Gravity. 
 
 Per cent. 
 H 2 SO 4 . 
 
 1 .000 
 
 o.oo 
 
 1.320 
 
 41.50 
 
 1 .640 
 
 71.99 
 
 I .010 
 
 i-57 
 
 1-33 
 
 42.66 
 
 .650 
 
 72.88 
 
 1.020 
 
 3-03 
 
 1.340 
 
 43-74 
 
 .660 
 
 73.64 
 
 1 .030 
 
 4-49 
 
 350 
 
 44.82 
 
 .670 
 
 74-51 
 
 1.040 
 
 5-9 6 
 
 .360 
 
 45.88 
 
 .680 
 
 75-42 
 
 1 .050 
 
 7-37 
 
 37 
 
 46.94 
 
 .690 
 
 76.30 
 
 1. 060 
 
 8.77 
 
 .380 
 
 48.00 
 
 .700 
 
 77.17 
 
 1 .070 
 
 10. 19 
 
 390 
 
 49.06 
 
 1 .710 
 
 78.01 
 
 1 .080 
 
 ii .60 
 
 .400 
 
 50.11 
 
 1 .720 
 
 78.92 
 
 1 .090 
 
 12.99 
 
 .410 
 
 5I-I5 
 
 1-730 
 
 79.80 
 
 I . TOO 
 
 14-35 
 
 .420 
 
 52-15 
 
 1.740 
 
 80.68 
 
 I . IIO 
 
 i5-7i 
 
 i-43 
 
 53-n 
 
 1-750 
 
 81-56 
 
 1 . 1 20 
 
 17 .01 
 
 1.440 
 
 54-07 
 
 1.760 
 
 82.44 
 
 1.130 
 
 18.31 
 
 i-45o 
 
 55-03 
 
 1.770 
 
 83-32 
 
 1 . 140 
 
 19.61 
 
 i .460 
 
 55-97 
 
 1.780 
 
 84.50 
 
 1.150 
 
 20.91 
 
 1.470 
 
 56.90 
 
 1.790 
 
 85.70 
 
 1 . 1 60 
 
 22 . 19 
 
 1.480 
 
 57-83 
 
 1 .800 
 
 86.90 
 
 1 . 170 
 
 23-47 
 
 1.490 
 
 58.74 
 
 1.810 
 
 88.30 
 
 1.180 
 
 24.76 
 
 1.500 
 
 59-70 
 
 1.820 
 
 90.05 
 
 1 . 190 
 
 26.04 
 
 i .510 
 
 60.65 
 
 1.825 
 
 91 .00 
 
 I . 200 
 
 27.32 
 
 1.520 
 
 61.59 
 
 1.830 
 
 92. 10 
 
 I . 210 
 
 28.58 
 
 i-53o 
 
 62.53 
 
 1-834 
 
 93-05 
 
 I . 220 
 
 29.84 
 
 1-540 
 
 63-43 
 
 1-837 
 
 94.20 
 
 1.230 
 
 31.11 
 
 1-550 
 
 64.26 
 
 1-839 
 
 95-0 
 
 I . 240 
 
 32.28 
 
 i .560 
 
 65.08 
 
 i .840 
 
 95.60 
 
 1.250 
 
 33-40 
 
 i-57o 
 
 65-90 
 
 1.841 
 
 97.00 
 
 I . 260 
 
 34-57 
 
 1.580 
 
 66.71 
 
 1.8415 
 
 98.20 
 
 I . 270 
 
 35-7i 
 
 1-590 
 
 67-59 
 
 1.841 
 
 98.7 
 
 1.280 
 
 36.87 
 
 i .600 
 
 68.51 
 
 i .840 
 
 99.20 
 
 I .290 
 
 38-03 
 
 i .610 
 
 69-43 
 
 1.839 
 
 99.70 
 
 1.300 
 
 39-19 
 
 i .620 
 
 70.32 
 
 
 
 I .310 
 
 40.35 
 
 1.630 
 
 71.16 
 
 
 
408 
 
 EXPERIMENTAL CHEMISTRY. 
 
 TABLE XXIX. Continued. 
 (b) Hydrochloric Acid.* 
 
 Specific 
 Gravity. 
 
 Per cent. 
 HC1. 
 
 Specific 
 Gravity. 
 
 Per cent. 
 HC1. 
 
 Specific 
 Gravity. 
 
 Per cent. 
 HC1. 
 
 I .OOO 
 
 0.16 
 
 1.070 
 
 14.17 
 
 . 140 
 
 27.66 
 
 1 .005 
 
 I-I5 
 
 1-075 
 
 15. 16 
 
 145 
 
 28.61 
 
 I. 010 
 
 2.14 
 
 1.080 
 
 16. 15 
 
 .150 
 
 29-57 
 
 1 .015 
 
 3.12 
 
 1.085 
 
 17-13 
 
 155 
 
 30.55 
 
 1 .020 
 
 4-13 
 
 .090 
 
 i8.ii 
 
 .160 
 
 3I-52 
 
 1.025 
 
 5-15 
 
 095 
 
 19.06 
 
 -165 
 
 32-49 
 
 1.030 
 
 6.15 
 
 . IOO 
 
 20. 01 
 
 .170 
 
 33-46 
 
 1-035 
 
 7-15 
 
 .105 
 
 20.97 
 
 I-I75 
 
 34-42 
 
 1 .040 
 
 8.16 
 
 .no 
 
 21 .92 
 
 1.180 
 
 35-39 
 
 1.045 
 
 9. 16 
 
 1.115 
 
 22.86 
 
 1.185 
 
 36-31 
 
 1.050 
 
 10. 17 
 
 I .120 
 
 23.82 
 
 I . 190 
 
 37-23 
 
 1-055 
 
 ii. 18 
 
 I.I25 
 
 24.78 
 
 I-I95 
 
 38.16 
 
 I .060 
 
 12 . IQ 
 
 I.I30 
 
 25-75 
 
 I . I2O 
 
 39-n 
 
 1 .065 
 
 13.19 
 
 I-I35 
 
 26.70 
 
 
 
 * Lunge and Marchlewski in Zeit. f. angew. Chem. 1891, 133. 
 (c) Nitric Acid. (Lunge and Rey.) 
 
 Specific 
 Gravity. 
 
 Per cent. 
 HNO 3 
 
 Specific 
 Gravity. 
 
 Per cent. 
 HNO 3 
 
 Specific 
 Gravity. 
 
 Per cent. 
 HNO 3 
 
 I .OO 
 
 o.oo 
 
 1.18 
 
 29.38 
 
 1-36 
 
 57-57 
 
 I .Ol 
 
 1.90 
 
 1.19 
 
 30.88 
 
 1-37 
 
 59-3'9 
 
 1 .02 
 
 3-70 
 
 1 .20 
 
 32-36 
 
 1.38 
 
 61 . 27 
 
 1.03 
 
 5-50 
 
 I .21 
 
 33-82 
 
 1.39 
 
 63-23 
 
 1 .04 
 
 7.26 
 
 I .22 
 
 35-28 
 
 1.40 
 
 65-30 
 
 1-05 
 
 8.99 
 
 1.2 3 
 
 .36.78 
 
 1 .41 
 
 67.50 
 
 1. 06 
 
 10.68 
 
 -24 
 
 38.29 
 
 1.42 
 
 69.80 
 
 1.07 
 
 12.33 
 
 25 
 
 39.82 
 
 1-43 
 
 72.17 
 
 1. 08 
 
 13-95 
 
 .26 
 
 41-34 
 
 44 
 
 74-68 
 
 1 .09 
 
 15-53 
 
 .27 
 
 42.87 
 
 45 
 
 77.28 
 
 I . IO 
 
 17.11 
 
 .28 
 
 44.41 
 
 .46 
 
 79.98 
 
 I . II 
 
 18.67 
 
 1.29 
 
 45-95 
 
 47 
 
 82.90 
 
 I .12 
 
 20.23 
 
 1.30 
 
 47-49 
 
 .48 
 
 86.05 
 
 I-I3 
 
 21.77 
 
 I -3 I 
 
 49.07 
 
 49 
 
 89.60 
 
 I.I4 
 
 23-31 
 
 1.32 
 
 50-71 
 
 1.50 
 
 94.09 
 
 I-I5 
 
 24.84 
 
 i-33 
 
 52-37 
 
 i-5i 
 
 98. 10 
 
 1.16 
 
 26.36 
 
 i-34 
 
 54-07 
 
 1-52 
 
 99.67 
 
 1.17 
 
 27.88 
 
 i-35 
 
 55-79 
 
 
 
APPENDIX II. 
 
 409 
 
 TABLE XXIX. Continued. 
 (d) Ammonium Hydroxide.* 
 
 Specific 
 Gravity. 
 
 Per cent. 
 NH 4 . 
 
 Specific 
 Gravity. 
 
 Per cent. 
 NH 3 . 
 
 Specific 
 Gravity. 
 
 Per cent. 
 NH 3 . 
 
 0.882 
 
 34-95 
 
 0.922 
 
 21 . 12 
 
 0.962 
 
 9-35 
 
 0.884 
 
 34.10 
 
 0.924 
 
 20.49 
 
 0.964 
 
 8.84 
 
 0.886 
 
 33- 2 5 
 
 0.926 
 
 19.87 
 
 0.966 
 
 8-33 
 
 0.888 
 
 32.50 
 
 0.928 
 
 19-25 
 
 0.968 
 
 7.82 
 
 0.890 
 
 31-75 
 
 0.930 
 
 18.64 
 
 0.970 
 
 7-31 
 
 0.892 
 
 3l.5 
 
 0.932 
 
 18.03 
 
 0.972 
 
 6.80 
 
 0.894 
 
 30.37 
 
 0-934 
 
 17.42 
 
 0.974 
 
 6.30 
 
 0.896 
 
 29.69 
 
 0.936 
 
 16.82 
 
 0.976 
 
 5-8o 
 
 0.898 
 
 29.01 
 
 0.938 
 
 l6.22 
 
 0.978 
 
 5-30 
 
 0.900 
 
 28-33 
 
 0.940 
 
 I5-63 
 
 0.780 
 
 4.80 
 
 0.902 
 
 27.65 
 
 0.942 
 
 15.04 
 
 0.982 
 
 4-3 
 
 0.904 
 
 26.98 
 
 0.944 
 
 14.46 
 
 0.984 
 
 3.80 
 
 0.906 
 
 26.31 
 
 0.946 
 
 13.88 
 
 0.986 
 
 3-30 
 
 0.908 
 
 25-65 
 
 0.948 
 
 I3-3I 
 
 0.988 
 
 2.80 
 
 0.910 
 
 24.99 
 
 0.950 
 
 12.74 
 
 0.990 
 
 2.31 
 
 0.912 
 
 24-33 
 
 0.952 
 
 12. 17 
 
 0.992 
 
 1.84 
 
 0.914 
 
 23.68 
 
 0-954 
 
 II .60 
 
 0-994 
 
 i-37 
 
 0.916 
 
 23-03 
 
 0.956 
 
 11.03 
 
 0.996 
 
 0.91 
 
 0.918 
 
 22.39 
 
 0.958 
 
 10.47 
 
 0.998 
 
 0-45 
 
 0.920 
 
 21-75 
 
 0.960 
 
 9.91 
 
 1 .000 
 
 o.oo 
 
 * Lunge and Wiernick in Ziet. f. angew. Chem. 1889, 183. 
 
4io 
 
 EXPERIMENTAL CHEMISTRY. 
 
 TABLE XXX. 
 
 Proportion by Weight of Absolute Alcohol. 
 (Mendeleef.)* 
 
 Specific gravity 
 atisC. 
 
 Per cent, 
 of alcohol. 
 
 Specific gravity 
 at 15 C. 
 
 Per cent 
 of 
 alcohol. 
 
 Specific 
 gravity 
 at 15 C. 
 
 Per cent, 
 of 
 alcohol. 
 
 0.9991 
 
 o-5 
 
 0.9501 
 
 34 
 
 0.8773 
 
 68 
 
 0.9981 
 
 I 
 
 0.9491 
 
 35 
 
 0.8750 
 
 69 
 
 0.9963 
 
 2 
 
 0-9473 
 
 36 
 
 0.8726 
 
 70 
 
 0.9945 
 
 3 
 
 0-9455 
 
 37 
 
 0.8702 
 
 7i 
 
 0.9928 
 
 4 
 
 0.9436 
 
 38 
 
 0.8678 
 
 72 
 
 0.9912 
 
 5 
 
 0.9417 
 
 39 
 
 0-8655 
 
 73 
 
 0.9896 
 
 6 
 
 0.9397 
 
 40 
 
 0.8631 74 
 
 0.9881 
 
 7 
 
 0.9377 
 
 4i 
 
 0.8607 75 
 
 0.9867 
 
 8 
 
 0-9357 
 
 42 
 
 0.8582 76 
 
 0.9853 
 
 9 
 
 0.9336 
 
 43 
 
 0.8558 
 
 77 
 
 0.9839 
 
 10 
 
 0.9316 
 
 44 
 
 0.8534 
 
 78 
 
 0.9826 
 
 ii 
 
 0.9294 
 
 45 
 
 0.8510 
 
 79 
 
 0.9813 
 
 12 
 
 0.9273 
 
 46 
 
 0.8485 
 
 80 
 
 0.9801 
 
 13 
 
 0.9251 
 
 47 
 
 o . 8460 
 
 81 
 
 0.9789 
 
 14 
 
 0.9230 
 
 48 
 
 0.8435 
 
 82 
 
 0.9777 
 
 15 
 
 0.9208 
 
 49 
 
 0.8410 
 
 83 
 
 0.9765 
 
 16 
 
 0.9186 
 
 50 
 
 0.8386 
 
 84 
 
 0-9753 
 
 i7 
 
 0.9164 
 
 5i 
 
 0.8360 
 
 85 
 
 0.9741 
 
 18 
 
 0.9142 
 
 52 
 
 0-8335 
 
 86 
 
 0.9728 
 
 i9 
 
 0.9119 
 
 53 
 
 0.8309 
 
 87 
 
 0.9716 
 
 20 
 
 0.9097 
 
 54 
 
 0.8283 
 
 88 
 
 0.9704 
 
 21 
 
 0.9074 
 
 55 
 
 0.8257 
 
 89 
 
 0.9691 22 
 
 0.9052 
 
 56 
 
 0.8230 
 
 90 
 
 0.9678 23 
 
 0.9029 
 
 57 
 
 0.8203 
 
 9i 
 
 0.9665 24 
 
 0.9097 
 
 58 
 
 0.8176 
 
 92 
 
 0.9651 25 
 
 0.8983 
 
 59 
 
 0.8149 
 
 93 
 
 0.9637 
 
 26 
 
 0.8960 
 
 60 
 
 0.8120 
 
 94 
 
 0.9623 27 
 
 0.8937 
 
 61 
 
 0.8092 
 
 95 
 
 0.9608 28 
 
 0.8914 
 
 62 
 
 0.8063 
 
 96 
 
 0-9593 2 9 
 
 0.8890 
 
 63 
 
 0.8034 
 
 97 
 
 0-9577 30 
 
 0.8867 
 
 64 
 
 o . 8004 
 
 98 
 
 0.9561 31 
 
 o . 8844 
 
 65 
 
 0.7973 
 
 99 
 
 0-9544 3 2 
 
 0.8820 
 
 66 
 
 0.7942 
 
 IOO 
 
 0.9527 
 
 33 
 
 0.8797 
 
 67 
 
 
 
 *Pogg. Ann. 138, p. 103. 
 
APPENDIX II. 
 
 411 
 
 TABLE XXXI. 
 
 Proportion by Volume of Absolute Alcohol. 
 (Mendelee/.)* 
 
 ioo volumes spirits. 
 
 ioo volumes spirits. 
 
 ioo volumes spirits. 
 
 Specific gravity 
 at 15.5 C. 
 
 Contain 
 volumes of 
 alcohol. 
 
 Specific gravity 
 at 15.5 C. 
 
 Contain 
 volumes 
 ot 
 alcohol 
 
 Specific 
 gravity 
 
 15-5 c. 
 
 Contain 
 volumes 
 of 
 alcohol. 
 
 I . OOOO 
 
 O 
 
 0.9604 
 
 34 
 
 0.8950 
 
 68 
 
 0.9985 
 
 I 
 
 0.9591 
 
 35 
 
 0.8925 
 
 69 
 
 0.9970 
 
 2 
 
 0-9577 
 
 36 
 
 0.8901 
 
 70 
 
 0.9956 
 
 3 
 
 0.9563 
 
 37 
 
 0.8876 
 
 7i 
 
 0.9942 
 
 4 
 
 0.9548 
 
 38 
 
 0.8851 
 
 72 
 
 0.9928 
 
 5 
 
 0-9534 
 
 39 
 
 0.8825 
 
 73 
 
 0.9915 
 
 6 
 
 0.9518 
 
 40 
 
 0.8800 
 
 74 
 
 0.9902 
 
 7 
 
 0.95 3 
 
 4i 
 
 0.8774 
 
 75 
 
 0.9889 
 
 8 
 
 0.9486 
 
 42 
 
 0.8747 
 
 76 
 
 0.9877 
 
 9 
 
 0.9470 
 
 43 
 
 0.8721 
 
 77 
 
 0.9866 
 
 10 
 
 0.9454 
 
 44 
 
 0.8694 
 
 78 
 
 0.9854 
 
 ii 
 
 0.9436 
 
 45 
 
 0.8667 
 
 79 
 
 0.9844 
 
 12 
 
 0.9419 
 
 46 
 
 0.8640 
 
 80 
 
 0.9832 13 
 
 o . 9400 
 
 47 
 
 0.8611 
 
 81 
 
 0.9822 
 
 14 
 
 0.9382 
 
 48 
 
 0.8583 
 
 82 
 
 0.9811 
 
 15 
 
 0.9364 
 
 49 
 
 0.8554 
 
 83 
 
 o . 9$q i 
 
 16 
 
 0.9344 
 
 50 
 
 0.8525 
 
 84 
 
 0.9790 
 
 17 
 
 0-9325 
 
 5i 
 
 0.8496 
 
 85 
 
 0.9781 
 
 18 
 
 o.93 5 
 
 52 
 
 0.8466 
 
 86 
 
 0.9771 19 
 
 0.9285 
 
 53 
 
 0.8435 
 
 87 
 
 0.9761 
 
 20 
 
 0.9265 
 
 54 
 
 o . 8404 
 
 88 
 
 0-9751 
 
 21 
 
 0.9244 
 
 55 
 
 0.8372 
 
 89 
 
 0.9741 
 
 22 
 
 0.9222 
 
 56 
 
 0.8340 
 
 90 
 
 o.973i 
 
 23 
 
 0.9201 
 
 57 
 
 0.8306 
 
 9i 
 
 0.9720 
 
 24 
 
 0.9180 
 
 58 
 
 0.8272 
 
 92 
 
 0.9709 
 
 25 
 
 0.9158 
 
 59 
 
 0.8236 
 
 93 
 
 0.9699 
 
 26 
 
 0.9139 
 
 60 
 
 0.8199 
 
 94 
 
 0.9688 
 
 27 
 
 0.9113 
 
 6! 
 
 0.8161 
 
 95 
 
 0.9677 
 
 28 
 
 o . 9090 
 
 62 
 
 0.8121 
 
 96 
 
 0.9667 
 
 29 
 
 0.9067 
 
 63 
 
 o . 8080 
 
 97 
 
 0.9654 
 
 30 
 
 0.9045 
 
 64 
 
 0.8035 
 
 98 
 
 0.9642 
 
 31 
 
 0.9022 
 
 65 
 
 0.7989 
 
 99 
 
 0.9630 
 
 32 
 
 0.8997 
 
 66 
 
 0.7939 
 
 IOO 
 
 0.9617 
 
 33 
 
 0.8974 
 
 67 
 
 
 
 *Pogg. Ann., 138, 230. 
 
412 
 
 EXPERIMENTAL CHEMISTRY. 
 
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 nd 
 r 
 y 
 
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 Mosander 
 Demarca 
 Mon 
 
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 F-H ^fH ^J 
 
 alas 
 
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APPENDIX II. 
 
 413 
 
 v>\O ^O n Tf jr< OO ON "r! f^- O ^ rt O u">OO M CNI rj- NI 
 ONO O M O .22 ON rO.22 i-i fO *^.22 ONOO ON 10 t^ O t- O 
 
 00 t^-00 00000 OOOOUOOOOt^O t^OO 00 t^ J>.00 t^OO 
 
 
 
 >-. 
 
 cfl 
 
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 PQ 
 
 ^ 
 
 M * 
 
 H 
 
 
 
 11 II : I 
 
 ,5 <4 o g . cj 
 
 tl 
 
 LOM eq OOOONOOsOOONONOrOOO 
 
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 H t^ ON MCMONiOOOrOO CNliOO ON-^-CNIOIHONIH 
 
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 O) 
 
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 I B 1 
 
 fc J2 "o i; ^ ^^o ^gg^rf^^ ogg.y.ss^^^^^ 
 
 OO O ffiffi M^^^SSS g^^^^OOfiH^d,^ 
 
414 
 
 EXPERIMENTAL CHEMISTRY. 
 
 LOW ro M to O\ O *~>- M ^ 
 
 oo o o o ^tt^t^M (N.JJ 
 
 CO OGO OO CO CO OO CO OO CJ 
 
 OOOO {J t^t^ 
 
 g 
 
 W 
 
 if 
 
 8-S 
 
 oi o 
 
 J 
 
 
 
 
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 :? 
 
 
REFERENCE BOOKS. 
 
 Dictionaries. 
 
 Comey, Arthur M. Dictionary of Chemical Solubilities. The Macmillan Co., 
 
 New York. $5.00. 
 Seidell, Atherton. Solubilities of Inorganic and Organic Substances. D. Van Nos- 
 
 trand Co., N. Y. 
 Thorpe. Dictionary of Applied Chemistry. Longmans, Green & Co., N. Y. 3 vols. 
 
 $50.00. 
 Watts, Henry. Dictionary of Chemistry. Revised by Morley and Muir. Longmans, 
 
 Green & Co., N. Y. 4 vols. $50.00. 
 
 Bibliographical and Historical Chemistry. 
 
 Alembric Club Reprints. The University of Chicago Press, Chicago. 40 cents each 
 
 Armitage, F. P. A History of Chemistry. Longmans, Green & Co., N. Y. $1.50. 
 
 Bolton. Bibliography of Chemistry. Chemical Pub. Co., Easton, Pa. $3.50. 
 
 Ladenburg, Dr. A. Lectures on the History of the Development of Chemistry Since 
 the Time of Lavoisier. The University of Chicago Press, Chicago. $1.75. 
 
 Meyer, Ernest von. History of Chemistry. Trans, by G. McGowan. The Mac- 
 millan Co., N. Y. $4.50. 
 
 Muir, M. M. Pattison. A History of Chemical Theories and Laws. Wiley & Sons, 
 N. Y. 
 
 Muir, M. M. Pattison. Heroes of Science, Chemists. Thomas Nelson & Sons. 
 $1.50. 
 
 Ramsey, W. Gases of the Atmosphere. The Macmillan Co., N. Y. $2.00. 
 
 Thorpe, Thomas E. Essays in Historical Chemistry. The Macmillan Co. $2.25. 
 
 Venable, Frank Preston. Short History of Chemistry. D. C. Heath & Co., % Boston. 
 $1.00. 
 
 Venable, F. A. Development of the Periodic Law. Chemical Publishing Co., Easton, 
 Pa. $2.50. 
 
 Descriptive and General Chemistry. 
 
 Bloxam. Inorganic and Organic Chemistry. P. Blakiston's Son & Co., Phila. 
 
 (New Ed.) 
 
 Gooch and Walker. Outlines of Inorganic Chemistry. The Macmillan Co. 
 Holleman. Inorganic Chemistry. Trans, by Cooper. John Wiley & Sons., N. Y. 
 
 $2.50. 
 
 Jones, H. C. Principles of Inorganic Chemistry. The Macmillan Co., N. Y. $4.00. 
 Mendeleeff, Dimitry Ivanovitch. Principles of Chemistry. Longmans, Green & 
 
 Co., N. Y. 2 vols. $10.00. 
 Newth, G. S. Text-book of Inorganic Chemistry. Longmans, Green & Co., N. Y. 
 
 $1.75. 
 Ostwald, Wilhelm. The Principles of Inorganic Chemistry. The Macmillan Co., 
 
 N. Y. $6.00. 
 
 Remsen, Ira. Chemistry (Adv. Course). Henry Holt & Co., N. Y. $2.80. 
 Richter, Victor von. Inorganic Chemistry. P. Blakiston's Son & Co., Phila. $1.75. 
 Roscoe and Schorlemmer. Treatise on Chemistry (Revised, 1905; i vol. ready). The 
 
 Macmillan Co., N. Y. 
 Smith, Alexander. General Inorganic Chemistry. The Century Pub. Co., N. Y. 
 
 415 
 
416 REFERENCE BOOKS. 
 
 Organic Chemistry. 
 
 Gattermann, Ludwig. Practical Methods of Organic Chemistry. Trans, by Shober. 
 
 The Macmillan Co , N. Y. $1.60. 
 Holleman, A. F. A Text-book of Organic Chemistry. Trans, by Walker. John 
 
 Wiley &Sons, N. Y. $2.50. 
 Lassar-Cohn. Laboratory Manual of Organic Chemistry. Trans, by Smith. The 
 
 Macmillan Co., N. Y. $2.25. 
 
 Mann, Gustav. Chemistry of the Proteids. The Macmillan Co., N. Y. 
 Noyes, W. A. Organic Chemistry. D. C. Heath & Co., Boston. 
 Orndorff, W. R. Laboratory Manual in Organic Chemistry. D. C. Heath & Co., 
 
 Boston. 35 cents. 
 
 Perkin and Kipping. Organic Chemistry. J. B. Lippincott Company, Phila. 2 vols. 
 Remsen, Ira. Organic Chemistry. D. C. Heath & Co., Boston. $1.20. 
 Richter, Victor von. Organic Chemistry. Trans, by Smith. P. Blakiston's Son & 
 
 Co. 2 vols. $6.00. 
 
 Theoretical and Physical Chemistry. 
 
 Arrhenius, Svante. Theories of Chemistry. Longmans, Green & Co., N. Y. 
 
 Ames. Scientific Memoirs. American Book Co., Cin. 
 
 Dobbin, L. and W. J. Chemical Theory for Beginners. The Macmillan Co., New 
 York. 70 cents. 
 
 Freund, I. The Study of Chemical Composition. The Cambridge Press, Cam- 
 bridge, Eng. 
 
 Findlay, Alexander. Practical Physical Chemistry. Longmans, Green & Co., N. Y. 
 
 Jones, Harry C. Elements of Physical Chemistry. Theory of Electrolytic Dis- 
 sociation. The Macmillan Co., N. Y. 
 
 Le Blanc, Max. Electro-Chemistry. Trans, by Whitney and Brown. The Mac- 
 millan Co. 
 
 Meyer, Oskar Emil. The Kinetic Theory of Gases. Trans, by Baynes. Longmans, 
 Green & Co., N. Y. 
 
 Meyer, Lothar. Outlines of Theoretical Chemistry. Trans, by Bedson and Williams. 
 Longmans, Green & Co., $2.50. 
 
 Morgan, J. L. Elements of Physical Chemistry. John Wiley & Sons., N. Y. 
 
 Nernst, W. Theoretical Chemistry. Trans, by R. A. Lehfeldt. The Macmillan 
 Co., N. Y. 4v,i 
 
 Remsen, Ira. Theoretical Chemistry. Lea Brothers & Co., Phila. $2.00. 
 
 Talbot and Blanchard. The Electrolytic Dissociation Theory with some of its Appli- 
 cations. The Macmillan Co., N. Y. 
 
 Walker, James. Introduction to Physical Chemistry. The Macmillan Co., N. Y. 
 $3.00. 
 
 Wurtz. The Atomic Theory. D. Apple ton & Co., N. Y. $1.50. 
 
 Physiological Chemistry. 
 
 Arrhenius, Svante. Immunochemistry. The Macmillan Co., N. Y. 
 
 Gamgee. Chemistry of the Animal Body. The Macmillan Co., N. Y. 2 vols. 
 
 $9.00. 
 Hammarsten, Olof. A Text-book of Physiological Chemistry. Trans, by Mandel. 
 
 John Wiley & Sons., N. Y. $4.00. 
 
 Industrial Chemistry. 
 
 Groves and Thorp. Chemical Technology. P. Blakiston's Son & Co., Phila. 
 
 4 vols. $16.00. 
 
 Thorpe, Frank H. Outlines of Industrial Chemistry. The Macmillan Co. $3.50. 
 Sadtler, S. P. Industrial Organic Chemistry. J. B. Lippincott Company, Phila. 
 
 $5.00. 
 
REFERENCE BOOKS. 417 
 
 Analytical Chemistry. 
 
 Bottger, Wilhelm. Qualitative Analysis. Trans, by Smeaton. P. Blakiston's Son 
 
 & Co, Phila. 
 
 Classen, Alexander. Quantitative Analysis. George Wahr, Ann Arbor, Mich. 
 Coblentz, Virgil. Volumetric Analysis. P. Blakiston's Son & Co., Phila. 
 Dennis, L. M. Hempel's Gas Analysis. The Macmillan Co, N. Y. 
 Fresenius. Quantitative Chemical Analysis. Trans, by Cohn. Wiley & Sons, 
 
 N. Y. 2 vols. $12.50. 
 
 Leffman and Beam. Food Analysis. P. Blakiston's Son & Co, Phila. 
 Newmann, Bernhard. Electrolytic Methods of Analysis. Trans, by Kershaw. 
 
 Whittaker & "Co, N. Y. 
 Olsen, J. C. Quantitative Chemical Analysis. D. Van Nostrand Company, N. Y. 
 
 $4.00. 
 
 Sulton, Francis. Volumetric Analysis. P. Blakiston's Son & Co, Phila. $5.00. 
 Stillman, Thomas B. Engineering Chemistry. Chemical Publishing Co, Easton, Pa. 
 Treadwell, F. P. Analytical Chemistry. Tran. by Hall. John Wiley & Sons, N. Y. 
 
 2 Vols. $7-00. 
 
 Miscellaneous. 
 
 Cornwall, H. B. Manual of Blowpipe Analysis. D. Van Nostrand Co, N. Y. 
 Dana, E. S. Minerals and How to Study Them. John Wiley & Sons, N. Y. $1.50. 
 Fournier, E. E. The Electron Theory. Longmans, Green & Co, N. Y. $1.50. 
 Hardin : W. L. Rise and Development of Liquefaction of Gases. The Macmillan 
 
 Co, $1.50. 
 
 McMillan, W. G. Electro-Metallurgy. Charles Griffin & Co, London. 
 Rutherford, E. Radio-Activity. University Press, Cambridge, Eng. 
 Shenstone, W. A. Methods of Glass-Blowing. Longmans, Green & Co, N. Y. 
 
 50 cents. 
 
 Tables. 
 
 Landolt and Bornstein. Physikalisch-chemische Tabellen. G. E. Stechert & Co, 
 
 N. Y. 
 Ostwald, Wilhelm. Physico-Chemical Measurements. Trans, by Walker. The 
 
 Macmillan Co, N. Y. 
 Traube, Dr. J. Physico-Chemical Methods. Trans, by Hardin. P. Blakiston's 
 
 Son & Co, Phila 
 
INDEX. 
 
 Acetamide, 248 
 Acetone, 248 
 Acetates, 259 
 
 solubility of, 260, 400 
 Acetylene, 234 
 Acid*, acetic, 248, 259 
 
 arsenic, 220 
 
 arsenious, 220 
 
 boric, 272 
 
 bromic, 160 
 
 carbolic, 256 
 
 carbonic, 230 
 
 chloric, 155 
 
 chloroplatinic, 368 
 
 chromic, 345, 346 
 
 dithionic, 190 
 
 formic, 248 
 
 fumaric, 250 
 
 hexathionic, 190 
 
 hydrazoic, 199 
 
 hydriodic, 162 
 
 hydrobromic, 157 
 
 hydrochloric, 107, 108, 409 
 
 hydrocyanic, 237 
 
 hydrofluoric, 152, 153 
 
 hydrofluosilicic, 153, 243 
 
 hypobromous, 159 
 
 hypochlorous, 154 
 
 hyposulphurous, 190 
 
 iodic, 163, 164 
 
 maleic, 250 
 
 manganic, 352 
 
 metaphosphoric, 217 
 
 metastannic, 335 
 
 molybdic, 349 
 
 muriatic, see Hydrochloric acid, 107, 
 1 08, 409. 
 
 nitric, 208, 209, 210, 408 
 oxidation by, 210 
 test for, 212 
 
 nitrosyl sulphuric, 186 
 
 nitrous, 212 
 
 test for, 212 
 
 oleic, 262 
 
 orthophosphoric, 216 
 
 oxalic, 260, 261 
 
 *A11 acids are listed under "Acids." 
 
 Acid, palmitic, 262 
 
 pentathionic, 190 
 
 perchloric, 155 
 
 perchromic, 346 
 
 periodic, 165 
 
 persulphuric, 190 
 
 phosphoric, 216 
 
 picric, 256 
 
 prussic, see Hydrocyanic acid, 237 
 
 pyrophosphoric, 217 
 
 pyrosulphuric, 190 
 
 pyrosulphurous, 190 
 
 selenic, 191 
 
 silicic, 241 
 
 stannic, 338 
 
 stearic, 262 
 
 sulphuric, 184 
 
 chamber process, 185 
 contact process, 188 
 dissociation of, 187 
 heat of formation of, 189 
 heat of solution of, 186 
 properties of, 186 
 
 sulphurous, 183 
 
 tartaric, 260 
 
 thiocyanic, 238 
 
 thiosulphuric, 189, 190 
 
 trithionic, 190 
 
 telluric, 191 
 
 tungstic, 350 
 Acidic salts, 116, 117 
 Acidimetry, nS 
 
 and alkalimetry, 285 
 Acids, 81, 107 
 
 activity ("strength") of, 141, 142 
 
 characteristic properties, no 
 
 "condensed," 272 
 
 definitions of, in, 117, 141 
 
 general properties of, no 
 
 ionization of, 145 
 
 monobasic, dibasic, etc., 117 
 
 nomenclature of, in 
 
 organic, in, 248, 259, 260 
 
 oxascids and hydrascids, in 
 Acids and bases, interaction of, 114 
 Acids, bases and salts, 104 
 
 419 
 
420 
 
 INDEX. 
 
 Affinity, chemical, 36, 51 
 Affinity-constant, 120 
 Air, see Atmosphere. 
 Albumin, 249 
 Alcohol, absolute, 253 
 
 ethyl, 252, 253 
 
 methyl, 248 
 
 table of, 420, 421 
 
 tests for, 254 
 Alcohols, 248 
 Aldehyde, 248 
 
 preparation of, 258 
 Alkalies, 114 
 Alkaline earth metals, 306 
 
 detection of, 314 
 Alkali metals, 280 
 
 detection of, 290 
 Alkalimetry and acidimetry, 285 
 Allotropy, definition, 38 
 Alloys, 276, 277 
 
 table of, 402, 403 
 Alum, preparation of common, 330 
 
 iron, 359 
 Alumina, 332 
 Aluminates, 329 
 Aluminum, 328, 329 
 
 properties and compounds, 329-331 
 Aluminum family, 271 
 
 comparative table, 275 
 Aluminum salts, analytical reactions of, 
 
 . 33i 
 
 Aluminothermy, 329 
 Alums, classification of, 331 
 Amalgams, 316 
 
 ammonium, 289 
 Amethyst, 242 
 Amides, 248 
 Amines, 248 
 Ammonia, 199 
 
 preparation and properties, 200 
 
 volumetric composition, 203 
 
 weight of a liter of, 202 
 Ammonium, 288 
 
 compounds, 201, 288 
 
 hydroxide, table, 409 
 
 salts, reactions of, 289 
 Analysis or decomposition (type of 
 
 chemical change), definition, 28 
 Analysis, qualitative, of anions, 371 
 
 cations, 370 
 
 Anhydride, meaning of term, in 
 Anhydrous, definition, 78 
 Anion, 126 
 
 Anions, names of, 135 
 Anode, 126 
 Antimony, 221 
 
 properties and compounds, 222-223 
 
 table, 225 
 
 test for, 222 
 
 Apatite, formula, 315 
 
 Aqua regia, 221 
 
 Aqueous (vapor) tension, 84 
 
 table of, 389, 390 
 Argentum, see Silver, 292, 299 
 Argon, 197 
 Arrhenius, 124, 127, 129 
 
 theory of electrolytic dissociation, 
 
 130, 131 
 Arsenates, 220 
 Arsenic, 218 
 
 properties and compounds, 221 
 
 table, 225 
 
 tests for 218, 225 
 Arsenites, 220 
 
 Arsenuretted hydrogen (arsine) 218 
 Arsine, 218 
 Asbestos, 314 
 
 platinized, 68 
 Atmosphere, 193 
 
 composition, etc, 193-198 
 
 table of, 394 
 Atom, definitions of, 50 
 Atomic, heats, table of, 395 
 
 hypothesis, 50 
 
 volume, 174, 175 
 
 weight, 50 
 
 weights, 412 
 Aurates, 304 
 Avogadro's rule, 49, 128 
 Azurite, 305 
 
 B 
 
 Balanced action, 122 
 Balance, chemical, 379 
 Bancroft, W. D., quoted, 37 
 Barite (heavy spar), 215 
 Barium, 313 
 
 properties and compounds, 313-314 
 
 tests for, 314 
 
 Barometer, Table of corrections for, 386 
 Baryta water, 231 
 Bases, 81, 113, 117, 142 
 
 activity (strength) of, 142 
 
 characteristic properties, 113 
 
 classification of, 113-114 
 
 general properties, 113 
 
 ionization, table of, 145 
 
 common ion, from, 142 
 
 monacid, diacid, 117 
 
 nomenclature, 114 
 
 "strong," 113 
 Basic salts, 116, 117 
 Basicity, 117 
 Basic oxides, 114 
 Bayley's table, 173 
 Bauxite, 332 
 Becquerel rays, 334, 350 
 Beilby, G. T., 133 
 
INDEX. 
 
 421 
 
 Bell-metal, see Table of Alloys, 402, 403 
 
 Bending glass, instructions, 375 
 
 Benzene, 246 
 
 Berlin (Prussian) blue, 360 
 
 Berthollet, on law of definite proportions, 
 
 45 
 
 Berthollet's laws of chemical reaction, 122 
 Berthelot's proposed third law of thermo- 
 chemistry, 39 
 
 Berzelins' "dualistic theory," 125 
 Bessemer process, see Iron, 356 
 Bibliography, 415-417 
 Bicarbonate of soda, 230 
 Binary compounds, 1 1 1 
 Bismuth, 223 
 
 properties and compounds, 223-224 
 
 table of properties, 225 
 Blast lamp, 374 
 Bleaching, by chlorine, 105 
 
 by sulphur dioxide, 182 
 Bleaching powder, 154, 310 
 Blue-stone (blue vitriol) see Copper sul- 
 phate 
 
 Blow-pipe, use of, 240 
 Blow-pipe flame, oxidation with, 240 
 
 reduction with, 240 
 Body, term defined, 18 
 Boiling-point, elevation of, 98 
 
 constants, 99 
 Boracite, 272 
 
 Borates, solubility of, 275, 400 
 Borax, 272 
 Borax bead, 273 
 
 colors of, 274 
 Boron, 271 
 
 properties and compounds, 271-274 
 
 table of properties, 275 
 
 test for, 274 
 Boyle's law, 20 
 Brass, see Alloys, 276, 277 
 Brimstone, see Sulphur, 176, 177 
 Britannia Metal, see Alloys, 276, 277 
 Bromides, solubility of, 159, 400 
 Bromine, 156 
 
 properties and compounds, 156-160 
 
 table of properties, 165 
 Bronze, 335 
 
 see Table of Alloys, 402, 403 
 Brookite, 333 
 Bunsen burner, study of, 338, 373 
 
 flame, 239 
 
 temperature of flame, 240 
 Burettes, 376-377 
 Butter, 262 
 
 Cadmium, 316 
 
 properties and compounds, 321-322 
 salts, analytical reactions of, 322 
 
 Caesium, 290 
 Calamine, 316 
 Calcite, 315 
 Calcium, 309 
 
 properties and compounds, 309-312 
 
 tests for, 311 
 
 Calculations in chemistry, 52 
 Calomel, 323 
 Calories, kinds of, 52 
 Camphor gum, 228 
 Cane-sugar, 265, 268 
 
 structural formula for, 251 
 Caramel, 268 
 Carbohydrates, 248 
 Carbon, 227 
 
 amorphous, 227 
 
 disulphide, 236 
 
 properties of, 227-229 
 
 halides of, 236 
 
 hydrides of, 233-235 
 
 oxides of, 229-232 
 Carbon and nitrogen, 236, 237 
 Carbon compounds, some common, 246- 
 
 270 
 Carbon family, 226 
 
 table of, 244 
 Carbonates, 230 
 
 solubility of, 231 
 
 test for, 230 
 Carborundum, 243 
 Carnallite, 308 
 Cassiterite, 335 
 
 Catalysis (catalytic action), 58, 69 
 Catalyzer, definition of, 58 
 
 action of, 58, 59 
 Cathode, 126 
 Cation, 126 
 Cations, names of, 135 
 Caustic, lunar, 303 
 Caustic alkalies, 114 
 Cavendish, 54 
 Cellulose, 265, 269 
 Cement, 315 
 
 hydraulic, 329 
 Cerite", 326 
 Cerium, 326, 333 
 Chalcopyrite, 356 
 Chalk, 315 
 
 Chalybeate water, 355 
 Chamber process, 185 
 Charcoal, 227, 228 
 Charles' law, 20 
 
 Chemical action, illustrations of, 22, 26, 
 27, 28, 29 
 
 means of initiating by heat, 26, 32, 
 36 
 
 by solution, 29 
 
 Chemical actions, reversible, 119 
 Chemical affinity, 36, 51 
 
422 
 
 INDEX. 
 
 Chemical changes, varieties of, 28, 29, 31, 
 
 33 
 
 characteristics of, 26-30 
 Chemical energy, 36, 51 
 Chemical equilbrium, 119, 121 
 
 laws of, 122 
 
 Chemical equivalent, 48 
 Chemical phenomena, 22 
 Chemical properties, 33, 54 
 Chemical reactions, complete, 119 
 
 incomplete, 119, 121 
 
 reversible, 119, 124 
 
 laws of, 122 
 Chemistry and physics, relation between, 
 
 17, 18 
 
 Chemistry, organic, 227 
 Chemistry, science of, 17 
 Chemistry, thermo-, 38, 39 
 China ware (porcelain), 327 
 Chlorates, 1-55 
 
 solubility of, 155 
 
 test for, 155 
 Chlorides, 164 
 
 solubility, 154 
 
 test for, 347 
 Chlorine, 104, 153 
 
 properties and compounds, 104-110, 
 
 I 53~ I 5 6 
 
 table of properties, 165 
 Chloroform, 261, 262 
 Chromates, 345, 346 
 
 oxidizing power of, 346 
 Chrome-alum, 347 
 Chrome-yellow, 341 
 Chromic, anhydride, 346 
 Chromites, 345 
 Chromium, 344 
 
 "passive", 344 
 
 properties and compounds, 344-349 
 
 salts, analytical reactions of, 348 
 Chromous compounds, 344 
 Chromyl chloride, 347 
 Cinnabar, 316, 323 
 Clay, 329 
 Coal, distillation of, 235 
 
 gas, 235 
 
 tar > 335 
 Cobalt, 355, 363 
 
 properties and compounds, 363-364 
 
 salts, reactions of, 364 
 Cobalt amines, 363 
 Cobaltite, 363 
 Cobaltous compounds, 363 
 Coke, 235 
 
 Colloidal solution, 83, 242 
 Columbium (niobium), 343 
 Combination (synthesis), 29 
 Combining proportions, laws of, 45, 46 
 
 volumes, law of, 48 
 
 Combining weights, law of, 48 
 Combustion of carbon in oxygen, heat of, 
 
 58 
 Common ion, effect of adding, ionic 
 
 equilibrium, 146 
 Complex ions, see Ions 
 
 salts, 144, 298 
 
 Component, definition of physical, 32 
 Compound, 32 
 
 endothermic, 39 
 
 exothermic, 39 
 Compounds, binary, in 
 
 ternary, in 
 Concentration, 90 
 
 expression of, 90 
 Concentration (mass) effect, 120 
 Concentration (ionic) effect, 320 
 Conductivity, electrical, 135-138 
 
 equivalent, 138 
 
 interpretation of, 131, 136, 137 
 Conservation of energy, 21 
 
 of mass, 1 8, 43 
 Constant, affinity, 120 
 
 electro-chemical, 134 
 
 ion-product, 146 
 
 ionization, 139 
 Constants, molecular depression, 100 
 
 molecular elevation, 99 
 
 physical, table of, 393 
 Constituent, definition of chemical, 32 
 Constitution of substances, 249 
 
 of alcohol, proof of, 250 
 Copper (cuprum), 292 
 
 analytical reactions of salts of, 297 
 
 equivalent of, 296 
 
 properties and compounds of, 292- 
 
 299 
 
 Copperas, 359 
 Copper group, 292 
 Corks, boring of, 375 
 
 rubber, treatment of, 376 
 Corpuscles, see Electrons, 132 
 Corrosive sublimate, 323 
 Corundum, 329 
 Cotton, 269 
 
 Critical phenomena, 49 
 Critical solution temperature, 95 
 Cry of tin, 335 
 Cryolite, 152 
 Cryp to-crystalline, 241 
 Crystallization, water of, 78, 79 
 
 see Hydration 
 
 Cupric compounds, see Copper 
 Cuprous compounds, see Copper 
 Curie, M. and Mme., on radioactive sub 
 
 stances, 350 
 Cyanates, 237 
 Cyanides, 237 
 Cyanogen, 236 
 
INDEX. 
 
 423 
 
 D 
 
 Dalton's law of partial pressures, 62 
 
 Davy, Sir Humphry, 112, 125 
 
 Decomposition (analysis), 28 
 heat of, 39 
 
 Decrepitation, 78 
 
 Definite proportions, law of, 45 
 
 Deflagration, 287 
 
 Degrees of ionization, table of, 145 
 
 Dehydration, 78 
 
 Deliquescence, 78 
 
 Density, 19 
 
 tables of, 384, 385 
 
 Depression constant, molecular, 100 
 
 Depression of the freezing-point, 97, 98, 
 100 
 
 "Developer" (in photography), 301 
 
 Dextrose, 270 
 
 Dialysis, 242 
 
 Diamond, 227 
 
 Dichromates, 345, 346 
 
 Didymium, 326 
 
 Dielectric constants, 132 
 
 Diffusion of gases, law of, 67 
 table of, 394 
 
 Diffusion phenomena in solution, 83 
 
 Dilution formula, Ostwald's, 138, 139 
 
 Displacement of ions, by a free metal, 146 
 by a free non-metal, 147 
 power of metals, order of, 147 
 
 Dissociation constant, 139, 142 
 
 Dissociation, electrolytic, 124-151 
 
 hydrolytic, see Hydrolysis, 77, 80, 144 
 in solution, proofs of, 124-151 
 of compounds by heat , 123 
 
 Distillation, 77 
 
 Dobereiner's triads, 168 
 
 Double decomposition (metathesis), 33 
 
 Double (compound) salts, 144, 298 
 
 Dulong and Petit's law, 52, 395 
 
 Dyeing, see Mordants, 331 
 
 Dynamic condition of chemical equilib- 
 rium, 123 
 
 Dynamite, 254 
 
 Dysprosium, 412 
 
 "Earth" metals, 306 
 Earths, alkaline, 306 
 Efflorescence, 78 
 Eka -aluminum, (gallium), 327 
 Eka-boron (scandium), 327 
 Electrical energy and chemical energy, 148 
 Electro-chemical equivalent, 134 
 Electrolysis, definition of, 126 
 Electrolysis and electrical equivalents, 134 
 Electrolyte, 126, 135 
 
 Electrolytic determination of copper in 
 cupric sulphate, 297 
 
 Electrolytic dissociation theory, 130 
 origin and development of, 124 
 
 Electromotive force (potential), 149 
 series, 150 
 
 E. M. F. as a measure of the tendency 
 toward chemical reaction, 149, 150 
 
 Electronic theory, 40, 132 
 
 Electrons, 132 
 
 Electro-plating, definition of, 364 
 
 Element, definition of, 37 
 
 Element versus "elementary substance," 
 
 37 
 
 Elementary (simple) substances, 31 
 Elements, acid-forming, 118, 152 
 
 base-forming, 118, 276 
 
 classification of, 167 
 
 grouping of, 279 
 
 metalic, 31, 118, 152 
 
 non-metallic, 31, 118, 276 
 
 table of, 412 
 Emery, 329 
 Enantiotrophy, 324 
 Endothermic compounds, 39 
 
 preparation of, 301 
 
 Energetics of chemistry, note on the, 34 
 Energy, 20 
 
 bound, 37 
 
 chemical, 36 
 
 conservation of, 21 
 
 definition of, 20 
 
 form, 24 
 
 free or available, 37 
 
 internal, 21, 35 
 
 forms of, 21, 35 
 
 transformation of, 21 
 "Energy content," 34 
 Energy and chemical changes, 29, 36 
 Energy and matter, relation of, 22, 23 
 English measures, 382 
 Equations, kinds of, 50 
 Equilibrium, 119 
 
 chemical, 119, 121, 122 
 
 ionic, 146 
 
 study of ionic, 320 
 Equivalents, definition of, 48 
 
 chemical, 48 
 
 electro-chemical, 134 
 Erbium, 326 
 
 Esters (ethereal salts), 248, 254 
 Ethers, 248, 257 
 
 compound, 258 
 Ethyl acetate, 248, 254, 262 
 Ethyl alcohol, 248 
 
 properties and preparation, 252, 253 
 
 table of proportion by weight and 
 volume of, 410, 411 
 
 tests for, 254 
 Ethylene, 233, 246 
 Ethylene bromide, 233 
 
424 
 
 INDEX. 
 
 Europium, 412 
 Exothermic compound, 39 
 Experiment, term denned, 40 
 Extracts, fluid, 253 
 
 Fact, term denned, 40 
 "Factors," term denned, 33 
 Faraday, 126 
 
 "Faraday" (96,580 coulombs of elec- 
 tricity), 134 
 Faraday's laws, 126 
 Fat, 262 
 Fatty acids, 254 
 Fehling's solution, 406 
 Feldspar, 241 
 Fermentation, 252 
 Ferrates, 355 
 Ferric alum, 359 
 Ferric compounds, see Iron 
 Ferricyanides, 237, 360 
 Ferrocyanides, 237, 360 
 Ferro-manganese, 351 
 Ferrous compounds, see Iron, 355 
 Fertilizers, 311 
 
 "Fixing," in photography, 301 
 Flame, study of a, 238 
 
 Bunsen, 238 
 
 oxidizing and reducing, 239 
 "Flashing-point," 251 
 Flash-light powder, 307 
 Flax, 269 
 
 Flowers of sulphur, 177 
 Flint, 241 
 
 Fluorides, solubility of, 153, 400 
 Fluorine, 152 
 
 preparation and properties of, 152-153 
 
 table of properties of, 165 
 Fluorite (fluor-spar) 152 
 Fluor-spar, 152 
 Flux, 277 
 Fools' gold, 356 
 Formaldehyde, 259 
 Formalin, 259 
 
 Formation, heat of, 39, 397, 398, 399 
 Formulae, kinds of, 51 
 
 dualistic, 51 
 
 empirical, 50, 249 
 
 graphic or structural, 52, 250 
 
 rational, 51, 250 
 
 space, 52 
 Franklinite, 356 
 
 Freezing-point, depression of, 100 
 Fructose, 268 
 Fusion of ice, heat of, 79 
 
 Gadolinite, 326 
 Gadolinium, 326 
 
 Galena, 336 
 
 Gallium, 326, 327, 328 
 
 Garnet, 245 
 
 Gas, coal, 234, 235 
 
 laughing, see Nitrous oxide, 205, 207 
 
 illuminating, 235 
 Gas law, 49 
 Gases, critical constants of, 393 
 
 diffusion velocities of, 394 
 
 heat of solution of, 93 
 solubility of, 96, 97 
 Gasoline, 251 
 
 Gay-Lussac's law, see Charles' law, 20, 49 
 Gay-Lussac's law of gaseous volumes, 48 
 Gelatin, formula of, 238 
 Gelatinoids, 249 
 Germanium, 334 
 
 table of properties of, 337 
 German silver, 364, 402 
 Glass, etching of, 153 
 Glass, uranium, 350 
 
 Glass vessels, correction factors for cali- 
 brating, 388 
 Glass working, 3 74-3 75 
 Glauber's salt, 290 
 Glucinum, 306 
 Glucose, 265, 267, 268 
 Glucosides, 249 
 Glycerides, 255, 262 
 Glycerin, 255 
 Gold, 304 
 
 properties and compounds, 304-305 
 Graham and Bunsen's law of effusion, 67 
 
 table, 394 
 
 Goldschmidt's process, 328, 329, 330 
 Gram-molecule (mole), a chemical unit 
 
 of weight, 52 
 
 Gram-molecular volume, 52 
 Grape-sugar, 265 
 Graphite, 227 
 Grotthus, quoted, 125 
 Green fire, 314 
 Greenockite, 316 
 Guldberg and Waage, law of mass action, 
 
 120 
 
 Gun-cotton, 266 
 Gun-metal, composition of, 402 
 Gunpowder, 286 
 Gypsum, 310 
 
 H 
 
 Halides, 152, 248 
 
 Halite, see Sodium chloride 
 
 Halogen family, 152 
 
 Halogens, table of relations of, 164, 165 
 
 Hardness (water), permanent, 312 
 
 temporary, 231, 310 
 Hardness, scale of, 404 
 Heat of combustion of carbon, 58, 230 
 
INDEX. 
 
 425 
 
 Heat, of combustion, 58 
 
 of decomposition, 39 
 
 of formation, 39, 207 
 
 tables of, 397-399 
 
 of fusion of ice, 79 
 
 of ionization, 151 
 
 of neutralization, 143 
 
 of reaction, 39, 207 
 
 of solution, 92, 93 
 
 of vaporization of water, 79 
 "Heat sum," law of the "constancy of 
 
 the" (law of Hess), 39 
 Heavy-spar, 315 
 Helium family, 197 
 Hematite, 356 
 Hemp, 269 
 Henry's law, 96 
 
 Hess, law of (second law of thermo- 
 chemistry), 39 
 
 Heterogeneous systems, 18, 25 
 Hittorf, quoted, 127 
 Homogeneous systems, 18, 25 
 Homologous series, 246 
 Hiibnerite, 349 
 Hulett, quoted, 84 
 Hydrates, 80 
 Hydraulic cement, 329 
 Hydrazine, 199 
 Hydrides, ex., 73, 247 
 Hydrion, properties of, 141 
 
 velocity (absolute), 135 
 Hydrocarbons, 246 
 
 derivatives of, 247, 248 
 Hydrogele, 242 
 Hydrogen, 64 
 
 arsenuretted, se Arsine, 218 
 
 ionic, 141 
 
 nascent (active), 70 
 
 preparation and properties, 64-73 
 Hydrogen bromide, see Acid, hydrobromic, 
 
 J 57 
 Hydrogen chloride, 108, 109 
 
 see Acid, hydrochloric, 107, 409 
 Hydrogen dioxide (peroxide), 70, 71 
 Hydrogen fluoride, 152, 153 
 Hydrogen halides, 154 
 Hydrogen iodide, 162 
 
 dissociation of, 123 
 Hydrogen peroxide, 70, 71 
 Hydrogen sulphate, 186 
 
 see Acid, sulphuric 
 Hydrogen sulphide, 178, 179, 180 
 
 see Acid, hydrosulphuric. 
 Hydrolysis, 77, 80, 81, 144 
 
 of salts, 77, 144 
 Hydrosole, 242 
 
 Hydroxides, definition of, 113 
 Hydroxide-ion (hydroxidion), properties 
 of, 144 
 
 Hydroxide-ion, velocity (migration) of, 
 
 J 35 
 
 Hydro xyl, 113 
 Hygroscopic, 196 
 "Hypo," 302 
 
 Hypochlorites, oxidizing power of, 154 
 Hypothesis, atomic, 40 
 
 Avogadro's (Rule), 49, 128 
 
 definition of, 41 
 
 ionic, 130 
 
 kinetic-molecular, 49 
 
 Prout's, 167 
 
 I 
 Ice, see Water 
 
 heat of fusion of, 79 
 Iceland spar, see Calcite, 315 
 Identity, definition of, 18 
 Illuminating-gas, 235 
 Indicators, 360 
 
 table of, 405 
 Indigo, 154 
 Indium, 328 
 Induction in chemical reaction, period of, 
 
 164 
 
 Inductive capacity, specific, 132 
 Ink, 359 
 
 sympathetic, 364 
 Internal energy, 21, 35 
 Internal rearrangement, a kind of chemi- 
 cal change, 31 
 
 Iodide (tri-odide) of nitrogen, 301, 302 
 Iodide, starch, 161 
 Iodine, 152, 160 
 
 preparation and properties, 160-164 
 
 table of properties of, 165 
 
 test for, 161 
 
 tincture of, 161 
 lodoform, 261 
 Ion, definition of, 126, 130 
 
 effect of adding a common, 146 
 
 product, 146 
 Ionic, concentration, 146 
 
 equilibrium, 146, 320 
 Ionic hypothesis, 124, 130 
 
 development of, 124-130 
 
 objections to, 124 
 
 support for, 134-151 
 Ionic substances, 131, 135 
 Ionization, 130, 131, 136 
 
 change in volume during, 143 
 
 constant, 139 
 
 factors which influence, 136-140 
 
 heat of, 151 
 
 per cent, or degrees of, table of, 145 
 
 repression of, 363 
 
 scheme of, 298 
 lonogens, 131 
 
 classification of, 131 
 
426 
 
 INDEX. 
 
 Inogens, charges on, number of, 130, 
 Ions, chemical conduct of, 141 
 
 color of, 140 
 
 complex, 144 
 
 displacement of, 146, 147 
 
 nomenclature of, 135 
 
 color of, 140 
 
 source of electrical charges of, 131- 
 
 J 34 
 
 speed of migration of, 135 
 Indium, 367 
 Iron (ferrum), 355 
 
 alum, 359 
 
 carbide, 356 
 
 cast, 356 
 
 compounds, 355-362 
 
 electrolytic, 356 
 
 kinds of, 356 
 
 ores, 356 
 
 passivity of, 357 
 
 Pig, 35 6 
 
 properties of, 356-358 
 
 pure, 356 
 
 pyrites, 356 
 
 salts, analytical reactions of, 360, 
 361 
 
 spiegel, 356 
 
 wrought, 356 
 Iron elements, 355 
 Isatin, formula, of 154 
 Isomerism, 249 
 Isomers, 249 
 
 Jasper, 241 
 Jones, H. C., 82 
 
 K 
 
 Kalion, migration velocity, 135 
 Kaolin, 245, 329 
 Kerosene, fire test of, 251 
 flashing-point of, 251 
 Ketones, 248 
 Kindling temperature, 56 
 Kinetic-molecular hypothesis, 49 
 Kohlrausch, 127 
 Krypton, see Helium family, 197 
 
 Lamp-black, 228 
 
 Lanthanum, 326 
 
 Laughing-gas, see Nitrous oxide, 205, 207 
 
 Laplace, quoted, 39 
 
 Lavoisier, quoted, 39 
 
 Law, Berthollet's, of chemical reaction, 122 
 
 Berthelot's, proposed for thermo- 
 chemistry, 39 
 
 Boyle's, 49 
 
 Charles', 49 
 
 Law, Dalton's, 62 
 
 definition of, 40 
 
 Dulong and Petit's, 395 
 
 Faraday's laws, 126 
 
 gas, 49 
 
 Gay-Lussac's, 48 
 
 Graham's, 67 
 
 Henry's, 96 
 
 Hess's, 39 
 
 Le Chatelier's (theorem), 37 
 
 mass, Guldberg, and Waage's, 120 
 
 of combining weights, 48 
 
 of concentration effect, see Mass law 
 
 of conservation of energy, 21 
 
 of constancy of composition, 45 
 
 of definite proportions, 45 
 
 of depression of freezing-point, 100, 
 101 
 
 of diffusion, 67 
 
 of dilution, 138 
 
 of distribution, 96 
 
 of energetics, first, 21 
 
 of gaseous volumes, 48 
 
 of Hess, 39 
 
 of ion-product, 139 
 
 of matter, 18 
 
 of multiple proportions, 46 
 
 of partial pressures, 62 
 
 of partition, see Law of distribution 
 96 
 
 of reciprocal proportions, 48 
 
 of thermochemistry, 39 
 
 of vapor pressure of solvents, 129 
 
 Ostwald's dilution, 138 
 
 periodic, 166 
 
 Raoult's, 100, 101, 129 
 
 Van't Hoff's, 87 
 Lead, 333, 334, 336 
 
 properties and compounds, 336-341 
 
 salts, analytical reactions of, 341 
 
 sugar of, 340 
 
 table of properties, 337 
 
 white, 341 
 Lead-tree, 339 
 Le Blanc's theory of the source of the 
 
 electric charges on the ions, 131 
 Levulose, (fructose), 268 
 Light, in chemical action, 29 
 
 see Photography, 301 
 Lime, 309 
 Lime-water, 309 
 Limonite, 356 
 Litharge, 340 
 Lithium, 281 
 
 Litmus, as an indicator, 405 
 Lodestone, see Magnetite 
 Lothar Meyer's table, 170 
 Luminosity, of Bunsen burner flame, 
 cause of, 239 
 
INDEX. 
 
 427 
 
 Lunar caustic, 303 
 
 Lyes, 330, see Alkalies, 114 
 
 M 
 
 Magnesia alba (hydrated basic corbonate 
 
 of magnesium), 308 
 "Magnesia mixture," 217 
 Magnesium, 306, 307 
 
 properties and compounds of, 307- 
 
 39 
 
 salts, analytical reactions of, 308 
 Magnesium nitride, 196, 212 
 Magnetite, 356 
 Malachite, 305 
 Manganates, 351, 352, 353 
 Manganese, 351 
 
 properties and compounds of, 351- 
 
 354 
 
 Manganese alloys, 351 
 Manganese bronze, 402 
 Manganese spar, 351 
 Manganic compounds, 351 
 Manganin, 351 
 Manganites, 351 
 Manganous compounds, 351 
 Marble, 309 
 Marsh-gas, 233 
 Marsh's test, 215, 216 
 Mass, action, 120, 121, 122, 222 
 
 conservation of, 18, 43 
 
 effect, 120 
 Matrix, 277 
 Matter, 17, 18 
 Mayer, Julius Robert, 21 
 Meerschaum, 314 
 Mendeleef's table, 169 
 Meniscus, 377 
 
 Mercuric oxide, heating of, 28 
 Mercuric chloramide, 324 
 Mercurous chloramide, 324 
 Mercury, 316, 317 
 
 densities of, table of, 385 
 
 properties and compounds, 322324 
 
 reactions of salts of, 324, 325 
 
 table of properties of, 317 
 Metal, definition of a, 118 
 Metallic elements, see Metals 
 Metals, (base-forming elements), 276 
 
 "alkali," 280 
 
 "alkaline earth," 306 
 
 chemical properties of the, 277 
 
 classification of the, 278 
 
 earth, 326 
 
 grouping for purposes of analysis of 
 the, 370 
 
 occurrence in nature of the, 277, 278 
 
 physical properties of the, 276 
 Metallurgy, 277 
 Metamers, 249 
 
 Metathesis, 33 
 Methane (marsh gas), 233 
 Methods of science, 40, 41 
 Methyl, 247 
 Methyl, orange, 405 
 Mica, 245 
 
 Microcosmic salt, 217 
 Migration of ions, 135 
 
 speed of, 135 
 
 Milk of lime, see Lime-water, 309 
 Mineral, definition of, 277 
 Minium, 340 
 Mixture, 31, 32 
 Mohr's salt, 359 
 Moissan, 152 
 Molar, solutions, 90 
 
 volume, 52 
 
 weight, 52 
 Molecular, equations, 50 
 
 formulae, 50, 51 
 
 hypothesis, 49 
 
 weight, 50 
 
 methods of determining the, 52 
 Molecule, definitions of, 49 
 Molecules, compound, 50 
 
 monatomic, 51 
 
 simple, 50 
 Molybdenum, 349 
 
 properties and compounds of, 349 
 Monazite, 333 
 Mordants, 331 
 Morse and Frazer, 87, 88 
 Mortar, 310 
 Multiple proportions, law of, 46 
 
 N 
 
 Naphtha (naptha), 251 
 Naphthalene, 256 
 Nascent condition, 70 
 Nascent hydrogen, 70 
 Natrion, 135 
 
 Natrium, see Sodium, in, 280 
 "Negative," in photography, 301 
 Neodymium, 326 
 Neon, see Helium family, 197 
 Nernst, 93, 132, 148 
 Nessler's reagent, 406 
 Neutralization, 114, 117, 143 
 
 complete, 117 
 
 experimental study of,- 114 
 
 heat of, 117, 143 
 
 partial, 117 
 
 theory of, 143 
 
 thermochemistry of, 143 
 
 volume change of, 143 
 New elements predicted by the periodic 
 
 system, 327 
 
 Newland's octaves, 168 
 Niccolite, 364 
 
428 
 
 INDEX. 
 
 Nickel, 335, 364 
 
 properties and compounds, 364-365 
 
 reactions of salts of, 365 
 Nickelion, 365 
 Nickel-steel, 365 
 Nilson and Cleve discovered eka-boron 
 
 (scandium), 327 
 Niobium (columbium), 343 
 Nitrates, 211 
 
 effect of heat on, 211 
 
 oxidizing power of, 211 
 
 solubility of, 213, 400, 401 
 
 tests for, 212 
 Nitric, anhydride, 206 
 
 oxide, 205 
 
 Nitride of magnesium, 196, 214 
 Nitrites, 212 
 
 solubility of, 213, 400 
 
 tests for, 213 
 Nitrogen, 192, 199 
 
 compounds, 199-213 
 
 preparation and properties, 192 
 
 table of properties, 225 
 
 thermochemistry of oxides of, 205- 
 
 208 
 
 Nitrogen family, 199 
 Nitrogen tri-iodide, 201 
 Nitroglycerin, 254 
 Nitrosyl chloride, 211 
 Nitrosyl-sulphuric acid, 186 
 Nitrous, anhydride, 206 
 
 oxide, 205, 207 
 Nomenclature, 49-52 
 
 of acids, bases and salts, in, 114, 117 
 118 
 
 of ionic theory, 126, 130, 134-151 
 Non-electrolytes, 126, 135 
 Non-metal, definition of, 118 
 Non-metallic elements, grouping for 
 
 purposes of analysis, 371 
 Non-metals, grouping of, 167 
 Normal, salts, 117 
 
 solutions, 90 
 Noyes, A. A., quoted, 20, 85 
 
 O 
 
 Occlusion, 68, 69 
 
 Octaves, law of, 168 
 
 Oil, illuminating, see Kerosene, 251 
 
 Oil of vitriol, see Acid, sulphuric, 184 
 
 Oils, 262 
 
 drying, 263 
 
 essential or volatile, 262 
 
 fatty or fixed, 262 
 
 non-drying, 263 
 
 semi-drying, 263 
 Olein, 262 
 Oleomargarine, 263 
 Olivine, 314 
 
 Onyx, 242 
 
 Opal, 242 
 
 Open-hearth process, 356 
 
 Ores, 277 
 
 Organic chemistry, 227 
 
 Orthite, 326 
 
 Orthoclase, 245 
 
 Osmium, 367 
 
 Osmosis, 85 
 
 Osmotic pressure, 85, 88 
 
 definition of, 85 
 Osmotic pressure and molecular weights, 
 
 87 
 Osmotic pressure of cane sugar, 86 
 
 effect of concentration, 86 
 
 effect of temperature, 86 
 Ostwald, quoted, 59, 69, 119 
 Oxalates, solubility of, 261, 400 
 Oxidation, 54, 56 
 
 defined in terms of the ionic hypoth- 
 esis, 211 
 Oxides, 54 
 
 basic, 114 
 
 classification of, 57 
 Oxidizing flame, 240 
 Oxygen, 54 
 
 nascent, 107 
 
 preparation and properties of, 54-59 
 Oxygen and hydrogen, heat energy pro- 
 duced when they combine, 73 
 Oxygen and ozone, different amounts of 
 
 energy in, 59 
 Oxygen family, 176, 191 
 Ozone, 59, 176 
 
 Palladium, 367 
 Palmitin, 262, 269 
 Paper, 269 
 
 parchment, 265 
 Paraffin, 251 
 Paris, plaster of, 310 
 Partial pressures, law of, 62 
 "Passive" condition of elementary sub- 
 stances, 344, 357 
 Perchlorates, 155 
 Periodic arrangement, 171, 172, 173 
 
 of Lothar Meyer, 170 
 
 of Mendeleef, 169 
 Periodic law, 166 
 
 Periods of induction in chemical action, 1 64 
 Permanganates, 351, 352 
 
 oxidizing power of, 353 
 Petit and Dulong's law, 395 
 Petroleum, 251 
 Pewter, 339, 402, 403 
 Pfeffer, 128 
 
 Pfeffer's measurement of osmotic pressure 
 86 
 
INDEX. 
 
 429 
 
 Phase, definition of term, 24, 25 
 Phase rule, 37 
 Phenol, 256 
 Phenolphthalei'n, 257 
 Phenolphthalein, as an indicator, 405 
 Phenols, 248 
 Phenomenon, 18 
 
 surface, 29, 83, 84 
 Phosphates, 216, 217 
 
 solubility of, 218, 400 
 
 tests for, 218 
 Phosphine, 214 
 Phosphorus, 199, 213 
 
 -bronze, 402 
 
 chemical relations, table of, 225 
 
 properties and compounds of, 213- 
 218 
 
 properties of red and yellow, 38 
 
 see Phosphates, 216, 217 
 Photography, 301-302 
 Physical constants, table of, 393 
 Physical and chemical changes, 26 
 Pig, iron, 356 
 Pipettes, 378 
 Pitchblende, 350 
 Plaster of Paris, 310 
 "Platinized" asbestos, 68 
 Platinum, 366, 367 
 
 catalytic action of, 69, 368 
 
 properties and compounds of, 367- 
 369 
 
 reactions of salts of, 369 
 
 "spongy," catalytic action of, 68 
 "Platinum black," 367, 369 
 Plumbago, 227 
 Plumbites, 334 
 
 Plumbum, see Lead, 333, 336 
 Polymerization of water, 80 
 Polymers, 249 
 Polysulphides, 181 
 Portland cement, 315 
 "Positive" -photography, 301 
 Potash, caustic, 114 
 
 yellow prussiate of, 237 
 Potassium, 280, 286 
 
 properties and compounds, 286-288 
 
 radical present in the salt or com- 
 pound sought for, 288 
 
 salts of, see under the acid 
 
 solubility of compounds of, 288 
 
 tests for, 288 
 Potential, 149 
 
 Potentials, table of single, 150 
 Powder, smokeless, 266 
 Praseodymium, 326 
 Precipitates, terms used in describing the 
 
 structure of, 282 
 
 Precipitation, ionic, theory and formu- 
 lation of, 146 
 
 Prediction, definition of, 41 
 Prediction of new elements, 327 
 Pressure, critical, 49 
 
 osmotic, 84, 85, 87, 88 
 
 partial, 62 
 
 solution, 84, 147 
 
 vapor, 62, 84 
 Priestly, 54 
 Product, concentration, 146 
 
 ion, 146 
 
 solubility, 146 
 "Products," chemical, 33 
 Propane, 246 
 Properties, body, 19 
 
 chemical, 18, 54 
 
 definition of, 18 
 
 general, 18, 19 
 
 physical, 18 
 
 reaction, 54 
 
 specific, 18, 19 
 
 Proust established law of definite pro- 
 portions, 45 
 Prout's hypothesis, 167 
 Prussian blue, 360 
 Pseudo-solutions, 83, 242 
 Purification, of sodium carbonate, 283 
 
 of sodium chloride, 283 
 Purple of Cassius, 305 
 Pyrites, 356 
 Pyrolusite, 350 
 
 Qualitative analysis, cations, 371 
 
 cations, 370 
 
 Quantitative relationships, 43 
 Quantivalence, see Valence 
 Quartz. 241 
 Quartz-glass, 241 
 Quicklime, 309 
 Quicksilver (mercury), 316 
 
 R 
 
 Radicals, acid, in 
 
 organic, 247 
 Radio-activity, 350 
 Radium, 307, 314 
 Ramsay, 193, 197 
 Raoult's laws, 100, 101, 129 
 Rare earth metals, 326 
 Reaction, 33 
 
 heat of, 39 
 Reactions, balanced, 122 
 
 complete, 119, 121, 122 
 
 incomplete, 119, 122 
 
 reversible, 119 
 Realgar, 218 
 Red fire, 312 
 
 Red heat, temperature corresponding to, 
 404 
 
430 
 
 INDEX. 
 
 Red lead, 340 
 
 Red phosphorus, 38, 213 
 
 Reducing flame, 240 
 
 Reduction, in terms of ionic hypothesis, 
 
 211 
 
 Reversible reactions, 119, 121, 122, 183 
 
 Rhodium, 366 
 
 Rinmann's green, 320 
 
 Richards, T. W., 20, 31, 39, 40 
 
 Rochelle salt, 260 
 
 Rock crystal, see Quartz, 241 
 
 Rock salt (halite), see Sodium chloride 
 
 Rose's fusible metal, 403 
 
 Rubidium, 290 
 
 Ruby, 329 
 
 Rule, Avogadro's, 49, 128 
 
 definition of, 40 
 Rust, 306 
 
 Rusting of iron, 356 
 Ruthenium, 366 
 Rutile, 333 
 
 Saccharose, 265, 268 
 Sal ammoniac, 289 
 Sal soda, 291 
 Salt, common, 283 
 
 definition of, 117, 143 
 
 electrolysis of a, 114 
 
 Glauber's, 290 
 
 Rochelle, 260 
 
 impurities of common, 283 
 "Salting out" of soap, 263 
 Salts, 114 
 
 acid, 116, 117 
 
 anhydrous, 78 
 
 basic, 1 1 6, 117 
 
 classification of, 116, 117 
 
 complex, 144, 298 
 
 compound, 144, 298 
 
 dehydrated, 78 
 
 double or compound, 144, 298 
 
 essential, 262 
 
 ethereal, 248, 262 
 
 fixed, 262 
 
 hydrolysis of, 144 
 
 ionization of, per cent, of, 145 
 
 nomenclature of, 118 
 
 normal, 116, 117 
 
 solubilities of, 400, 401 
 Samarium, 326 
 Sand, 153, 241 
 Saponification, 264 
 Sapphire, 329 
 Saturated solution, 83, 84, 85, 90 
 
 making of a, 83 
 
 theory of, 83, 84 
 
 Saturation, a term used in organic 
 chemistry, 247 
 
 Scandium, 326 
 
 Scheele, 54 
 
 Science, definition of, 41 
 
 development of a, 40, 41 
 
 methods of, 41 
 
 natural, 41 
 Sciences, abstract, 42 
 
 biological, 41, 42 
 
 classification of, 41, 42 
 
 concrete, 42 
 
 descriptive or special, 42 
 
 general or speculative, 42 
 
 physical, 41 
 
 Scientific versus systematic method, 41 
 Selenite, 310 
 Selenium, 190 
 Series, electromotive, 150 
 
 in organic chemistry, 246 
 
 in periodic system, 169, 171 
 
 solution tension, 147 
 Serpentine, 245 
 Shot, making of, 336 
 Shot-metal, 336 
 Siderite, 356 
 
 Siemens-Martin process, 356 
 Silica, 153, 240 
 Silicates, 242, 243 
 
 solubility of, 243 
 
 test for, 243 
 Silicon, 226, 240 
 
 compounds of, 242, 243 
 
 preparation and properties of, 241 
 
 table of properties, 244 
 Silver (argentum), 292, 299 
 
 argenticyanide, 303 
 
 compounds, of 299-304 
 
 electrolytic deposition of, 299 
 
 preparation from an alloy of pure, 
 300 
 
 properties of, 299 
 
 reactions of salts of, analytical, 
 
 34 
 
 sterling, 300, 402 
 
 "tree," 299 
 Slag, 277 
 
 Slaked lime, 309, 310 
 Smaltite, 363 
 Smithsonite, 316 
 Soap, definition of a, 264 
 
 hard, 263, 264 
 
 preparation of, 263 
 
 properties of, 264 
 
 salting out of, 263 
 
 soda, 263 
 
 soft, 264 
 
 Soap-bubble solutions, 405, 406 
 Soda, ash, 282 
 Soda, caustic, 114 
 
 washing, 291 
 
INDEX. 
 
 431 
 
 Sodium (natrium), in, 280, 281 
 
 aluminate, 330 
 
 cobaltic nitrite, 285 
 
 hydroxide, 113, 282 
 
 properties of, 112, 113, 281 
 
 solubility of salts of, 285 
 
 salts and other compounds, 282-286 
 
 see also under the acid radical 
 present in the salt or compound 
 sought for. 
 
 zincate, 319 
 Solder, soft, 335 
 Solubilities, tables of, 400, 401 
 Solubility, data, 400, 401 
 
 effect of pressure on, 93, 96 
 
 effect of solvent on, 96 
 
 effect of surface on, 83, 84 
 
 effect of temperature on, 88, 89, 95. 97 
 
 explanation of degrees of solubility, 
 82 
 
 expression of, 90 
 
 measurement of, 89 
 
 mutual, 94 
 
 of gases, 96, 97 
 
 of liquids, 94 
 
 of pairs of liquids, 94 
 
 of solids, 82, 88, 89, 90, 91, 93 
 
 product, 146 
 Solute, definition of, 82 
 
 distribution of solute between two 
 
 immiscible solvents, 95 
 Solution, characteristics of a, 82 
 
 colloidal, 83, 221, 242 
 
 definition of a, 83 
 
 heat of, 92, 93 
 
 in two solvents, 95, 96 
 
 modern theory of, 124 
 
 partial explanation of the mechanism 
 of, 83, 84 
 
 process of, chemical or physical, 102 
 
 temperature, critical, 95 
 
 tension, 84, 146, 147, 148, 149 
 
 tension of metals, 146, 147 
 Solutions, 82 
 
 boiling-points of, 98 
 
 concentrated, 82, 90 
 
 dilute, 82 
 
 freezing-points of, 98-101 
 
 ionic, equilibrium of, 146 
 
 kinds of, 82 
 
 lowering of the vapor tension of, 97 
 
 molar, 90 
 
 normal, 90 
 
 pseudo-(colloidal), 83, 242 
 
 saturated, 82, 88 
 
 standard, 90 
 
 supersaturated, 89 
 " terminology of, 90 
 
 vapor tension of, 97 
 
 Solution tension, electrolytic, 149 
 
 hypothesis of, 84 
 
 lowering of the, 101 
 
 of the metals, 147 
 
 series, 150 
 
 Solvent, definition of, 82 
 Solvents, immiscible, 103 
 Soot, 228 
 Space formulae, 52 
 Specific gravities, 19, 173, 175 
 Specific heats of elements, 395 
 Specific inductive capacity of water, 132 
 
 of various solvents, 132 
 Specific volume, 19 
 Spectroscope, figures of, 280, 281 
 Speed of reaction affected by, catalyzers, 
 58, 68, 69, 1 88 
 
 concentration, 120, 146, 222, 320 
 
 solution, 29, 141 
 
 surface effects, 29, 66, 69, 307 
 
 temperature, 26, 28, 32, 33, 123, 188 
 Sphalerite, 316 
 Spiegel iron, 356 
 Spirits, 253 
 
 "Sponge," platinum, 68 
 Stannates, 334 
 Stannic chloride, 338 
 Stannic, compounds, see Tin, 333, 336 
 Stannum, see Tin, 333, 336 
 Starch, 266 
 
 test for, 266 
 
 theory of formation of, 269 
 Starch iodide, 161 
 Steam, see Water 
 Stearin, 262 
 Steel, 356 
 
 nickel, 365 
 Stibine, 222 
 Strontianite, 315 
 Strontium, 306, 312 
 
 analytical reactions of salts of, 312, 
 
 313 
 
 compounds, 312-313 
 
 ion, 313 
 
 properties of, 312 
 
 Structural or graphic formulae, 52, 250 
 Structure, molecular, 249, 250 
 
 see Constitution of substances, 249 
 Sublimate, corrosive, 323 
 
 see Mercuric chloride 
 Sublimation of iodine, 160 
 
 of salts of mercury, 323 
 Substance, compound, definition of, 32 
 
 definition of, 18 
 
 simple or elementary, definition of, 
 
 3*> 37 
 Substitution, 236 
 
 products, 236 
 Sugar, barley, 268 
 
432 
 
 INDEX. 
 
 Sugar, beet-, 268 
 
 cane-, 265, 268 
 
 grape-, 265 
 
 inversion of cane-, 265, 268 
 
 Fehling's test for, 265 
 
 invert, 268 
 
 see Saccharose 
 
 structural formula for cane-, 251 
 Sugars, classification of, 266 
 Sugar of lead, 340 
 
 see Lead acetate 
 
 Sulphates, solubility of, 187, 188 
 Sulphides, solubility of, 178, 180 
 
 test for, 1 80 
 Sulphites, solubility of, 183 
 
 test for, 183 
 Sulphur, 176, 177 
 
 allotropic forms of, 177, 178 
 
 amorphous, 178 
 
 chemical relations of, 178-190 
 
 compared with oxygen, 178 
 
 family, relationships of members of, 
 191 
 
 flowers of, 177 
 
 liquid forms of, 177 
 
 monoclinic, 177, 178 
 
 native, 178 
 
 oxygen derivatives of, 181 
 
 plastic, 177 
 
 properties of, 177, 178 
 
 rhombic, 177, 178 
 
 roll, 177, 178 
 
 transition points of, 177, 178 
 Sulphureted hydrogen, see Hydrogen 
 
 sulphide, 178, 180 
 Supercooled liquids, 177 
 Superphosphate of lime, 311 
 Surface phenomenon, 29, 66, 68, 69 
 Sylvite, see Potassium chloride 
 Symbols, 50 
 Sympathetic ink, 364 
 Synthesis, 29 
 
 Tables, list of, 
 Tantalum, 343 
 Tartar-emetic, 222, 260 
 Tartrates, solubility of, 260 
 Tellurium, 190 
 Temperature, critical, 199 
 
 critical solution, 95 
 
 kindling, 56 
 
 scale according to color, 404 
 Test paper, preparation of, 59, 106 
 Thallium, 328, 332 
 
 properties and compounds of, 332 
 Theory, definition of, 41 
 
 electronic, 132 
 
 of ions, 130 
 
 Thermit, 329 
 Thermochemistry, 38 
 
 laws of, 39 
 
 principles of, 39 
 
 Thermometer, corrections for, 383 
 Thermometric readings, conversion of, 
 
 383 
 
 Thiocyanates, 237 
 Thomas-Gilchrist process, 356 
 Thomson, J. J., 132 
 Thorium, 333 
 Tin (stannum), 333, 334, 335, 336 
 
 analytical reactions of salts of, 338 
 
 properties and compounds, 337-339 
 
 table of properties of, 337 
 Tincal (crude borax), 272 
 Tin cry, 335 
 Tincture, 253 
 Tin-plate, 336 
 Tin-stone (cassiterite), 335 
 Titanium, 333 
 Titration, 114, 115 
 Triads of Dobereiner, 168 
 Tridymite, 241 
 Triolein, 262, 264 
 Tripalmitin, 262, 264 
 Tristearin, 262, 264 
 Tungsten, 349 
 
 -steel, 350 
 
 Turnbull's blue, 360 
 Turpentine, 106 
 Type-metal, 336 
 
 composition of, 403 
 
 U 
 
 Ultramarine, 330 
 Units, chemical, 52 
 
 of electrochemistry, 134 
 Univalent, 41 
 
 Unsaturated compounds, 41 
 Uraninite, 350 
 Uranium, 350 
 Urea, 249 
 
 Valence (quanti valence), definitions of, 
 
 5 J > I 3 I 
 
 identical with electrical charge on the 
 ion, 131, 345 
 
 multiple, 51 
 
 table, 174 
 Vanadates, 343 
 Vanadium, 472 
 Van't Hoff, quoted, 128, 129 
 Vapor pressure, 62, 84 
 Vapor tension, lowering of the, 97 
 
 of liquids, 84 
 
 of water, tables of, 389, 390 
 
INDEX. 
 
 433 
 
 Vaporization of water, heat of, 79 
 
 Vapor pressure of mercury, 391 
 
 Vaseline, 251 
 
 Verdigris, 294 
 
 Vermillion, 223 
 
 Vinegar, see Acid, acetic, 248, 252 
 
 Vitriol, blue, 294, see Copper sulphate 
 
 green, 359 
 
 oil of, see Acid, sulphuric, 184 
 
 white, 319, 359 
 Volume, atomic, 174, 175 
 
 critical, 49 
 
 molar, 52 
 
 W 
 
 Waage and Guldberg's law of mass 
 
 action, 120 
 
 Walker, James, on osmotic pressure, 88 
 Washing soda, 291 
 Water, 75 
 
 as a solvent, 81 
 
 boiling temperature of, 392 
 
 catalytic properties of, 105 
 
 chemical properties of, 79 
 
 composition of, 75 
 
 critical temperature of, 79 
 
 density of, table, 384 
 
 dielectric constant of, 132 
 
 dissociation of, 79 
 
 electrolytic dissociation of, 79 
 
 electrolysis of, 27 
 
 formula of, variation of, 80 
 
 gas, 245 
 
 "hardness" of, 312 
 
 hydrolytic action of, see Hydrolysis, 
 77, 80, 144 
 
 of hydra tion (crystallization), 78 
 
 per cent, ionized, 131, 145 
 
 physical properties of, 79 
 
 polymerized, 80 
 
 purification of, 77 
 
 vapor tension of, tables, 389, 390 
 
 volume of one gram of, table, 385 
 Waters, chalybeate, 355 
 Water bottle, construction of, 376 
 Water-glass, 242 
 
 "Weak" acids and bases, explanation of 
 141-143 
 
 neutralization of, 143 
 Weighings in air reduced to vacuum, 
 
 table, 387 
 Weight, atomic, 50 
 
 molecular, 50 
 Weights, atomic, 412-414 
 
 chemically equivalent, 48 
 
 combining, 48 
 
 equivalent, 48 
 
 molecular, see under each metal 
 Welsbach mantles, 333 
 White lead, 341 
 
 chemical principles involved in the 
 
 manufacture of, 341 
 Wines, bouquet of, 258 
 Witherite, 315 
 Wohler, 249 
 Wolframite, 349 
 Wood, 269 
 
 Wood's metal, 220, 277 
 Wrought iron, 356 
 Wulfenite, 349 
 
 X 
 
 Xenon, 197 
 
 Yeast, 252 
 
 Yellow prussiate of potash, 237 
 Ytterbium, 326, 414 
 Yttrium, 326, 414 
 
 Zinc, 316, 317, 318 
 
 analytical reactions of salts of, 320 
 blende, 316 
 
 hydrogen equivalent of, 71 
 properties and compounds of, 318- 
 
 321 
 
 table of properties, 317, 318 
 vapor density determinations of, 317 
 
 Zincates, 317 
 
 Zircon, 333, 414 
 
 Zirconium, 333, 414 
 
 28 
 
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