Name 
 
 LABORATORY EXERCISES 
 
 CHEMISTRY 
 
 EDWIN J. BARTLETT. 
 
 Chapter 1. Object Lessons for Beginners. 
 
 Chapter 11. The Ordinary CalcuhUions. 
 
 Chapter 111. Experiments in Weij^ht and Volume Relations, 
 
 Chapter IV. Suggestions Hew to Make Chemicals. 
 
 Chapter V. Words and tlieir Uses. 
 
 ►f. .J- .j< 
 
 LEACH, SHEWELL, & SANBORN, 
 
 BOSTON. NEW YORK. CHICAGO. 
 
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 THE GIFT OF 
 
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Cornell University LIbrsry 
 arV19213 
 
 Laboratory exercises in chemistr 
 
 3 1924 031 262 896 
 oiin,anx 
 
Cornell University 
 Library 
 
 The original of tliis book is in 
 tlie Cornell University Library. 
 
 There are no known copyright restrictions in 
 the United States on the use of the text. 
 
 http://www.archive.org/details/cu31924031262896 
 
LABORATORY EXERCISES 
 
 IN 
 
 CHEMISTRY 
 
 BY 
 
 EDWIN J. BARTLETT 
 Professor of ChemistriFin Dartmouth College 
 
 LEACH, SHEWELL, & SANBORN, 
 
 BOSTON, NEW YORK, CHICAGO. 
 
Copyright, iSgi;, 
 By Edwin J. Bartlett. 
 
 PRESS OF SAMUEL U'^HEK, 
 BOSTON, MASS. 
 
TABLE OF CONTENTS. 
 
 PAGE 
 
 Preface 5 
 
 Introduction 7 
 
 Supplies 8 
 
 Apparatus 9 
 
 Chemicals 10 
 
 Chemical Notation 13 
 
 Chemical Names 15 
 
 Exercises. Part I 19 
 
 List for Identification 57 
 
 Chemical Computations. Part II 60 
 
 Table of Atomic Weights 61 
 
 Special Examples 70 
 
 Quantitative Relations. Part III 73 
 
 Chemical Preparations. Part IV 93 
 
 Chemical Terms. Part V 104 
 
 Reference Diagram 109 
 
 Index ..... .. iii 
 
PREFACE. 
 
 This book is a Laboratory Guide for beginners in chemistry, and 
 something more. 
 
 Part I is the common ground over which beginners must travel. 
 It has been many times worked over with reference to clearness of 
 instructions ; concentration of attention upon the essential thing by 
 avoiding complicated and distracting apparatus ; the elimination of 
 dangerous or seriously irritating experiments; economy in chemi- 
 cals ; relation to subsequent work. 
 
 It is thought especially adapted to conditions which are quite 
 prevalent — the necessity of, economy in supplies and teaching 
 force. The numerous references are essential and call the pupil's 
 attention to intersection of properties. 
 
 Part II contains the arithmetical guide, absolutely necessary in 
 some form for everything beyond the most elementary stage. It 
 tries to draw a careful distinction between those computations which 
 are original in their nature and those which merely utilize accumu- 
 lated facts. 
 
 If the pupil can solve the limited number of diverse illustrative 
 problems given, he will probably have no difficulty in the practical 
 cases which will arise. If the teacher prefers, he can easily multiply 
 the simpler problems. It should be remembered that here arith- 
 metic is a means, not an end. 
 
 Part III contains a series of quantitative experiments, which 
 illustrate fundamental principles. They are fairly accurate and do 
 not require complicated or costly apparatus, nor anything outside of 
 the rather familiar forms. It has been thought best to give full 
 working directions, but to leave explanations for the pupil's thought 
 with the guidance of the teacher. 
 
 Part IV is simple enough to those who have learned what it 
 
O PREFACE. 
 
 contains ; the instructions are merely those which every novice in 
 chemistry must have before he can pass to the slightest degree of 
 independence. Perhaps this Part will not be used as a distinct 
 course of study in many schools, but rather as containing convenient 
 directions to be utilized from time to time. 
 
 Part V contains a list, not exactly of definitions, but of terms and 
 their uses. They are not to be memorized, but learned in the 
 proper applications. 
 
 A large portion of Part II is essential to III ; but after the arith- 
 metic is applied Part III may be combined with 1; though it 
 seems to the writer preferable to have phenomena and manipulation 
 somewhat familiar first. 
 
 The time in the Laboratory (of course approximate) is Part I, 
 without bracketed experiments, 30 hours ; bracketed experiments, 10 
 hours ; Part III, 30 hours ; Part IV, 10 or more hours. 
 
 It will be seen that this book contemplates giving the beginner 
 the absolute essentials, — language, notation, substances, processes, 
 qualities, quantities, — all from the experimental standpoint. The 
 desirable but not immediately essential matters, — history, sources, 
 manufactures, applications, etc., — can easily be given by oral, text, 
 or reference instruction according to time available and according 
 to the pupil's interest. 
 
 The Atomic Theory should be treated when it can be understood 
 as an explanation of- facts and laws, not as a dogma to which" all 
 facts and laws must correspond. It is a well-fitting glove on a 
 substantial hand, but is not the hand. 
 
 A list of the chemicals and apparatus needed by the pupil for 
 the work laid out in the book is given ; but all the things desirable 
 and even necessary in a laboratory are not given. 
 
 The full list of materials for each day's work the teacher will find 
 in the paragraphs for the day, but for his convenience in ordering 
 goods and preparing reagents a list is given in order on p. 10. 
 This list does not repeat articles once provided for. 
 
INTRODUCTION. 
 
 The laboratory study of chemistry involves the study of 
 substances, of processes, of weight and energy relations. 
 The first two are for the last, and are in their nature 
 precedent. 
 
 The absolute essentials for the pupil's success are neat- 
 ness, orderliness, common sense, the ability to form a plan 
 from printed instructions and to adapt means to ends in car- 
 rying it out. 
 
 The pupil should always call things by their right names ; 
 and think of them, at any rate, by their distinguishing names. 
 There is a whole world of difference between saying and 
 thinking " that thing " and saying and thinking " a 6ocm. 
 royal Berlin porcelain evaporating dish, round bottomed, 
 lipped." 
 
 The duplication and repetition of experiments have great 
 value. By duplication the student tests his own work and 
 learns, too, of its regular, and not accidental, character ; by 
 repetition he corrects his mistakes of unfamiliarity, and uses 
 his own mind for improvement. 
 
 It is indispensable that scientific notes should be taken at 
 the time of the observations ; it is desirable that they should 
 be put into good form later. 
 
 All chemical properties are correlative. Zn-|-H2S04=Zn 
 SO^-j-Hg. It is equally a property of zinc to decompose 
 
 7 
 
8 INTRODUCTION. 
 
 sulfuric acid, and of sulfuric acid to give up its hydrogen 
 for zinc. 
 
 Every chemical action is dependent upon a certain set of 
 conditions. Success is a mere matter of acquired skill 
 within narrow lines. The interpretation and correction of 
 failure is far broader. 
 
 Processes in chemistry are direct inferences from the 
 properties, chemical or physical, of the substances involved. 
 
 The older chemist enlarges his knowledge and increases 
 his skill 
 
 1. By his own investigations, often — perhaps most often 
 — upon seeming trifles. 
 
 2. By reference to the great stock of accumulated knowl- 
 edge — text -books and monographs. 
 
 3. By noting the fresh results of others' work — in the 
 journals. 
 
 4. By taking numerical constants, except a few familiar 
 ones, from their special tables when he wants them. 
 
 The beginner should use the same methods. 
 
 Supplies. 
 
 In purchasing supplies for the laboratory it is to be taken 
 for granted that the work will continue, and material should 
 be obtained with reference to permanency of stock. It is 
 usually very poor economy to purchase little scraps in ounce 
 or half-ounce bottles. 
 
 Most of the chemicals, with reasonable care, last until they 
 are used. A few need special care. Strong .nitric acid 
 should be kept from the light ; lime, lime water, the strong 
 alkalis, all deliquescent and efflorescent salts, bisulfid of 
 
INTRODUCTION. 9 
 
 carbon, iodin, bromin, and all volatile substances must be 
 kept in tightly closed bottles. The concentrated acids are 
 purchased and diluted for use. Salts are purchased and dis- 
 solved in water as needed ; ordinary phosphorus is always 
 kept in water. Sodium and potassium are kept in naphtha, 
 or even kerosene. 
 
 Many of the articles in the following list are usually 
 obtainable near by, as the teacher will see upon looking it 
 over. Sugar, starch, lime, threepenny nails. No. 28 iron 
 wire, salt, bleaching powder, washing soda, etc., are usually 
 not far away. 
 
 As a source of heat, alcohol lamps are satisfactory except 
 in a few cases, as bending glass. When Bunsen burners are 
 used for heat, they must be watched very closely, as there is 
 a great tendency in the novice to push a burner under some- 
 thing and let it go, while he proceeds to consider something 
 in another part of the room. 
 
 It is both impracticable and undesirable to furnish each 
 pupil with any considerable individual stock of chemicals for 
 this particular work. The chemicals are best arranged for 
 the daily work in definite . order upon a common table. 
 Dilute acids and ammonia may be on the pupil's table. 
 
 The Apparatus, 
 
 The plan of these exercises contemplates the following 
 outfit for each pupil, which the teacher can modify to suit 
 himself in accordance with his experience and preference. 
 This is not intended to be a complete stock for a laboratory. 
 
 1 test tube rack. 
 
 12 test tubes 4" or 6"x}i". 
 
 2 beakers, plain, about 50 and 7Sc.c. 
 I evaporating dish, about 6oc.c. 
 
lO 
 
 INTRODUCTION. 
 
 2 side-necked test tubes, 6"xi". 
 
 I flat-bottomed Erlenmeyer flask, 200-250C.C. 
 
 1 thistle tube. 
 
 2 gas cylinders, 300C.C., ground edges. 
 
 2 cylinder covers, plates of glass 3" square cut from scrap window glass. 
 
 I funnel 2^". i clay crucible, 2 oz. 
 
 I piece of sheet iron, or iron netting 3" square. 
 
 I piece galvanized or tinned iron 6"x^", bent to a V for jar support. 
 
 I glass rod and spatula, 5" long, pointed at one end, flat at the other, made from 
 X" or i\" tubing. 
 
 I delivery tube, 15" of ^" tubing. 
 
 I low 2 qt. pan of earthenware or tin. 
 
 I iron ring stand. 
 
 I pair test tube holders. 
 
 I alcohol lamp, 4 oz., globe-shaped. 
 
 Rubber connections, corks to fit test tubes, side-necked test tubes and flask, 
 glass tubing as required 1^", J^", ^", ^"- 
 
 There will be needed occasionally in the pupil's work 50 or looc.c. graduates, 
 triangular files, Bunsen burners, cork borers, mortars and pestles, chemical ther- 
 mometers, scales, blast-lamp. 
 
 The materials needed in Parts III and IV are given there. 
 
 The chemicals for Part I arranged in order of use by para- 
 graphs. Those once given are not mentioned a second time. 
 The proportions are satisfactory for the purposes of this 
 book ; not for all other purposes. 
 
 2. scrap glass tubing. 
 
 3. crystallized copper sulfate, 
 ferric oxid. 
 
 4. dil. hydrochloric acid (1:3). 
 dil. ammonia (l : 2). 
 
 litmus paper, red and blue, cut in 
 
 narrow strips ij4" long. 
 3 or 4 other convenient reagents. 
 
 5. broken roll sulfur, 
 strips of copper foil, 
 dil. nitric acid (l : 3). 
 3-penny nails. 
 
 mercuric oxid. 
 
 splinters of straight-grained back- 
 ing. 
 
 zinc dust. 
 
 flowers of sulfur. 
 
 crystallized barium chlorid. 
 
 granulated zinc, made by melting 
 common scrap in an iron ladle 
 and pouring slowly into water. 
 
 cone, commercial sulphuric acid. 
 
 cone, commercial hydrochloric 
 acid. 
 
INTRODUCTION. 
 
 II 
 
 8. potassium chlorate, fine, ground, 
 
 if necessary, in a clean mortar. 
 
 sol. of silver nitrate (l : 40 distilled 
 water). 
 
 distilled water. 
 
 manganese dioxid, powdered. 
 
 iron wire (on spools. No. 28). 
 
 barium peroxid. 
 
 dil. starch paste. Shake two or 
 three lumps of common starch 
 with loc.c. of cold water, boil 
 till clear, add to about 20oc.c. 
 of water. 
 
 sol. of iodid of potassium (l : 40). 
 
 9. granulated sugar, 
 cryst. carbonate of soda, 
 quicklime in small pieces (par- 
 tially air-slaked will do). 
 
 lead oxid (litharge), 
 arsenious oxid (" white arsenic ") . 
 "fused" or anhydrous calcic 
 chlorid. 
 
 10. calcium sulfate (plaster of Paris), 
 hyposulfite of soda. 
 
 common table or dairy salt, 
 ordinary saltpeter. 
 
 11. hydrogen peroxid (ord. " 12 to 15 
 
 vol.") may be diluted with an 
 equal part of water. 
 
 sol. of potassium permanganate 
 (1:50). 
 
 fresh lead sulfid, made by 
 precipitating lead nitrate or 
 acetate with either hydrogen 
 sulfid or ammonium sulfid and 
 washing 3 or 4 times by decan- 
 tation. 
 
 12. paraffin (cut in lumps about 
 
 chestnut size), 
 hydrofluoric acid (in ceresine 
 bottle). 
 
 scrap window glass (cut about 
 
 2"x3"). 
 filter paper, common, 9.c.m to 
 1 1 cm. 
 
 13. "chlorid of lime" or bleaching 
 
 powder, obtainable in paste- 
 board boxes. 
 
 printed paper. 
 
 stale urine or other decomposing 
 material diluted with water, or 
 HjS water. 
 
 14. rock salt, broken. 
 
 strong ammonia (with great cau- 
 tion, for this experiment only). 
 
 15. small or broken crystals of potas- 
 
 sium bromid. 
 bromin, about J^ an inch in a 
 
 bottle, in the hood, must be 
 
 taken with a dropping tube and 
 
 not poured, 
 ferrous sulfid, broken into little 
 
 lumps. 
 
 16. small or broken crystals of potas- 
 
 sium iodid. 
 iodin crystals, 
 common alcohol, 
 bisulfid of carbon, 
 sol. of KI (same as 8). 
 baking powder. 
 
 17. 2 oz. hessian crucible. 
 
 ferrous sulfid (powdered and sifted) . 
 " dry " carbonate of soda, 
 selected charcoal (about egg size) . 
 
 18. lead nitrate or acetate sol. (i : 40 
 
 made clear by a few drops of 
 
 nitric or acetic acid), 
 antimony chlorid (l : 40 water 
 
 acidulated with HCl). 
 ferrous chlorid (i : 40, it may be 
 
 made by dissolving a nail in 
 
 cone. HCl). 
 
12 
 
 INTRODUCTION. 
 
 19. sheet copper (cut in strips about 
 
 extract of logwood (enough added 
 to water to give a good color) . 
 
 permanganate of n diluted with 
 3 or 4 parts of water. 
 
 sulfite of soda. 
 
 20. sol. of barium chlorid (i : 20). 
 sol. of sodium sulfate (i : 20). 
 sol. of di-sodium phosphate (l : 
 
 20). 
 
 21. ordinary ammonium chlorid. 
 commercial potassium nitrite. 
 
 22. ordinary phosphorus (cut into 
 
 pieces under water, and kept 
 
 under water till used), 
 lime water, see Part IV. 
 common cotton batting, 
 any of the commercial extracts of 
 
 beef, )4 teaspoonful to I qt. of 
 
 water. 
 
 23. sol. of ammonium chlorid or am- 
 
 monium sulfate (about 1 : 20) . 
 NaOH sol. (1 : 10) 
 
 24. sols, of stannous and cadmium 
 
 chlorids (1 : 40, fresh, may be 
 made by dissolving the metals 
 in cone. HCI). 
 ammonium sulBd made by satu- 
 raiing dil. ammonia with HjS 
 gas and then adding as much 
 more ammonia. 
 
 25. ammonium nitrate cryst. 
 
 26. nitric acid, sp. gr. 1.2. 
 
 28. stick tin is melted in a. Bunsen 
 
 flame, the drops fall upon a 
 smooth surface like slate, spread 
 out thin and are cut in strips, 
 [horse hair.] 
 
 29. [red or amorphous phosphorus.] 
 
 30. [calcium phosphid.] 
 
 31. sol. of CuSO, I: 20. 
 hypophosphorous acid(sp.gr.i.I5; 
 
 dilute with an equal volume of 
 water). 
 
 32. [glacial phosphoric acid.] 
 
 33. charcoal powdered and sifted. 
 
 34. metallic antimony broken in mor- 
 
 tar., » 
 
 natural sulfid of antimony, pow- 
 dered. 
 
 35. platinum, or iron wire, 
 borax " dry." 
 
 36. silicate of soda, soluble glass 
 
 (the commercial article in two 
 or three times its volume of 
 water) . 
 
 [fine sea sand. 
 
 powdered fluor spar. 
 
 quicksilver.] 
 
 37. powdered bone black. 
 
 38. formic acid (sp.gr. 1.06 will be 
 
 suitable). 
 
 39. carbonate of lime (marble dust), 
 turpentine. 
 
 40. caoutchouc (crude rubber gum, in 
 
 lumps as large as a pea). 
 
 41. lycopodium. 
 
 42. magnesium ribbon, 3". 
 cadmium (dropped like tin -28). 
 copper carbonate. 
 
 mercuric nitrate, 
 ferrous oxalate, 
 ferric chlorid sol. 1-20. 
 46. lead carbonate, 
 [ferrous chlorid.] 
 
 48. ammonium carbonate sol. i-io. 
 
 49. sodium carbonate sol. 1-20. 
 
 50. sol. of mercuric chlorid 1-40. 
 sol. of copper nitrate 1-20. 
 strips cut from sheet lead, 
 potassium ferricyanid sol. i : 20. 
 
INTRODUCTION. 1 3 
 
 Chemical Notation. 
 
 The chemical notation is merely a convenient system of 
 abbreviations expressing substances of chemical purity and 
 some essential facts about them. 
 
 The symbol of an element is an abbreviation of its Latin 
 name, using the initial letter, or the initial and some charac- 
 teristic subsequent letter. See Table, p. 6i. 
 
 Generally the Latin initial agrees with the English. The 
 following are the more important exceptions : — 
 
 Na. (Natrium) for Sodium. Pb. (Plumbum) for Lead. 
 
 K. (Kalium) for Potassium. Pe. (Ferrum) for Iron. 
 
 Sb. (Stibium) for Antimony. Hg. (Hydrargyrum) for Mercury. 
 
 Ag. (Argentum) for Silver. Sn. (Stannum) for Tin. 
 
 Au. (Aurum) for Gold. Cu. (Cuprum) for Copper. 
 
 In all accurate chemical notation the symbol means a cer- 
 tain definite proportion of the element called its atomic 
 weight. Thus O means i6 parts of oxygen and O^ 32 parts. 
 
 The symbol is, however, often employed to denote the sub- 
 stance, e. g. "Put a glowing splinter into the jar of O." 
 
 A formula represents the composition of a chemical com- 
 pound. It indicates what elements make up the compound 
 and the proportion of each element. Thus NaCl means a sub- 
 stance consisting of 23 parts of sodium and 35.4 of chlorin. 
 HNaCOg consists of i part hydrogen, 23 parts sodium, 12 
 parts carbon, and 48 oxygen. The whole is 84 ; so the ele- 
 ments are '+'3+"+4« . 
 
 In the case of gases and volatile liquids the formula repre- 
 sents the weight of definite standard volumes, usually called 
 "two volumes," or "the molecular volume." Thus HCl, 
 HjS, HgN, H4C — all gases — would represent equal volumes 
 in the same discussion. 
 
 In writing a formula the order of the elements is usually 
 
14 INTRODUCTION. 
 
 not important — KOH and KHO are the same; but in com- 
 pounds of two elements it is customary to write the more 
 metallic first — NaCl, not ClNa ; and in compounds of three 
 elements the usual order is metallic, non-metallic, oxygen — 
 KNO3, not NKO3, or O3NK. 
 
 A chemical formula implies that the composition of the 
 substance has been carefully determined by some one. For 
 the methods by which formulas are determined see Part II. 
 
 A chemical equation is a statement of a chemical change, 
 including all the substances that enter into and come from 
 the change in their relative proportions by weight; also, when 
 they are gases, in their relative proportions by measure. 
 
 Zn-j-2HCl=ZnCl2-|-H2 means that 65 parts of zinc dis- 
 place the hydrogen from 72.8 parts of hydrochloric acid, giv- 
 ing 135.8 of zinc chlorid and 2 of hydrogen. These numbers 
 are relative, but attaching a definite meaning to one fixes it 
 for all ; if 65 grammes, ounces, or pounds of zinc are used, 
 the other numbers represent the same denomination. The 
 numbers come from the Table of Atomic Weights, p. 61, and 
 methods of forming and using equations are given in Part II. 
 
 The pupil should remember that a chemical equation must 
 have been worked out experimentally by some one, or else it 
 is only a theory. Many of the equations in the books are 
 theories. 
 
 Many well established equations are given through Part I, 
 in such form that the pupil can complete them correctly. 
 
 The pupil should also remember that if an equation is 
 true, it is true for all the substances in it. Thus the equa- 
 tion given above may teach the behavior of zinc (with HCl), 
 or of HCl (with zinc), or it may apply to the making of H or 
 of zinc chlorid. 
 
 Familiarity with chemical formulas and equations is much 
 
INTRODUCTION. 1 5 
 
 like the spelling and meaning of words — to be acquired by 
 deliberate study, by constant use, and by looking up special 
 cases as they occur. While indispensable to sound knowl- 
 edge of chemistry, the mere knowledge of these symbols is 
 of no value as an end in itself. 
 
 The Chemical Notation is the means of expression of two 
 somewhat distinct ideas. 
 
 1. Weight and volume relations. 
 
 2. The conceptions of the Atomic Theory. 
 
 1. Certain facts, the result of experiment, which may be 
 
 looked upon as independent of any theory, find expression in 
 
 chemical formulas and equations. Thus NaCl expresses 
 
 what may be considered permanent facts, viz. that salt is 
 
 ■4^, sodium and-^* chlorin. 
 50.4 58.4 
 
 The mode of expressing these facts might at some time 
 change. Na=23, because that is. now thought upon sound 
 evidence to be the comparative weight of its chemical unit. 
 If at some time it should appear that the chemical unit of 
 sodium is 11.5, then the composition of salt would be ex- 
 pressed by NagCl. The composition would not change, but 
 the mode of expression would. 
 
 2. The prevailing and seemingly well-founded views of 
 matter known as the Atomic Theory are also carried into the 
 formulas, and we frequently say : " NaCl means a molecule of 
 the substance called salt, consisting of i atom of sodium and 
 I atom of chlorin." The evidence upon which this explana- 
 tion of chemical facts is based must come later. 
 
 Chemical Names. 
 
 Chemistry employs a system of names largely its own. 
 Each Element has a name based on : the old common name, 
 e. g. ferrum, iron ; some property, e. g. iodin, violet-like ; 
 
1 6 INTRODUCTION. 
 
 on some place or thing related, e. g. scandium, Scandinavia ; 
 calcium, calx, lime; on some myth, e. g. Thorium, Thor; 
 cerium, Ceres ; niobium, Niobe ; on some whimsical name, 
 e. g. nickel, kupfer-nickel, Old Nick's copper, etc. 
 
 Each Compound has a name which is an attempt to express 
 its composition and which, therefore, implies in some way 
 the names of the elements making it and their proportions. 
 Compounds of two elements have a name ending in id {ide). 
 NaCl — sodium chlorid, the more metal-like or electro-positive 
 element being named first and id being affixed to the less 
 metal-like. 
 
 When the same pair of elements enter into combination in 
 two different proportions (as SOj and SO3) the terminations 
 ous and ic are often employed ; both are sulfur oxids ; the first 
 svlUvlvous oxid, the second sulfur?'<? oxid. Numerals are also 
 conveniently used, sulfur dioxxA, sulfur trioxid. 
 
 In compounds of three elements ite and ate are commonly 
 used ; KNOj, potassium nitrite ; KNO3, potassium nitrate. 
 In a few compounds this system is carried farther by the use 
 of hypo and per as prefixes ; hypo . . . ite, less oxygen than 
 ite, per . . . ate, more oxygen than ate ; (KCIO, potassium 
 hypochlorite ; KCIO2, potassium chlorite ; KCIO3, potassium 
 chlorate ; KCIO^, potassium perchlorate.) See Part V. 
 
 Classes of compounds are grouped under common names. 
 Thus oxids, chlorids, sulfids, etc., are compounds of two ele- 
 ments, one of which is oxygen, chlorin, sulfur, respectively. 
 Nitrates are compounds containing the group NO3 (KNO3, 
 NaN03, Cu(N03)2, etc.). Sulfates contain the group SO^ 
 (Na2S04, CaSO^, Al2(S04)3, etc.). 
 
 Processes have names, e. g. Filtration, Evaporation, Dis- 
 tillation, etc. 
 
 Pieces of Apparatus have names, and in the ordinary labora- 
 
INTRODUCTION. 1 7 
 
 tory supplies are very accurately designated either by maker's 
 number, or better by dimensions or contents. Some good 
 apparatus catalogue is very useful in the laboratory. 
 
 To the Pupil. 
 
 Think. 
 
 Read each paragraph through before you begin it; and 
 form a plan. 
 
 Wear an apron or other suitable protection to the clothes. 
 
 Arrange your apparatus in a symmetrical, orderly way, and 
 keep it so. 
 
 Have plenty of cleaning cloths. 
 
 Put no liquids on anything too hot to hold in the hand. 
 
 When mixing strange chemicals make it a habit to be cau- 
 tious ; add a drop first. 
 
 Use your good sense. 
 
 Keep your notes parallel with your work, and look up all 
 the references. 
 
 Learn formulas as you would words, in spelling and mean- 
 ing. Do not think there is any sound rule for "writing 
 formulas " ; you will be much helped after a time by analogy 
 and valence ; but formulas are facts, not guesses. 
 
 You cannot make a chemical equation till you are well 
 advanced in knowledge and experimental skill ; but you must 
 begin at once to read. The equations given are meant for 
 reading lessons, with a little call upon your attention to 
 supply missing parts. 
 
 Learn the language. 
 
 Find out what is in this book and then go beyond it for 
 wider information. 
 
Missing Page 
 
LABORATORY EXERCISES IN 
 CHEMISTRY. 
 
 PART I. 
 
 LABORATORY EXERCISES. 
 
 Note. — Experiments in brackets should always be omitted until the instructor 
 has considered their suitability to his classes, laboratory, and other conditions. 
 
 Preliminary Manipulations, 
 
 1. Measure in cubic centimeters, and record the capacity 
 of beakers, test tubes, evaporating dish, etc. Consider 
 whether it is best to fill the beaker and pour into the gradu- 
 ated cylinder, or the reverse. 
 
 3. Cut, bend to a right angle, draw out, and close pieces 
 of glass tubing. This should be demonstrated by the 
 teacher, and practised upon scrap glass. 
 
 3. Solution, Suspension, Filtration, Evaporation. Half-fill 
 two test tubes with water ; into one put a crystal of sulfate 
 
 19 
 
20 EXERCISES IN CHEMISTRY. 
 
 of copper ; into the other about the same bulk of ferric oxid. 
 Shake and warm each tube. One substance finally dissolves, 
 giving a blue liquid ; the other gives a turbid mixture. Now 
 pour the contents of both tubes into a beaker together. 
 Prepare a filter (demonstrate), and pour all upon it, catching 
 the liquid which runs through — called filtrate — in evaporat- 
 ing dish. Evaporate the contents of the dish nearly, not 
 quite, to dryness, by placing it over the lamp upon the ring 
 stand. How have you separated water, oxid of iron, and 
 sulfate of copper .' 
 
 4. Testing reaction, i. Half -fill two test tubes with water; 
 into one put 2 or 3 drops of an acid (e. g. vinegar, acetic 
 acid), into the other put 2 or 3 drops of an alkali (e. g. 
 common ammonia, ammonic hydrate) ; into each tube dip 
 a strip of red litmus paper and a strip of blue litmus paper ; 
 note results ; now exchange the papers ; note results. An 
 acid and an alkali have what effect .'' 2. Now take any 
 four liquids in your reagent bottles and classify them into 
 acid, alkaline, and neutral by means of litmus paper. 
 
 Physical arid Chemical Properties and Action. 
 
 5. I. Examine a piece of sulfur and a strip of copper 
 foil ; write all the difference in properties you can discover 
 by examination alone. You will note in this way only phys- 
 ical properties. 2, Now cautiously heat a little of the sul- 
 fur in a test tube until it boils, and then, while it is boiling, 
 drop half the strip of copper to the bottom of the tube. 
 Keep it boiling a minute longer. Now examine the strip ; 
 how has it changed ? What is it ? 3. Put the rest of the 
 copper into dilute nitric acid in a test tube and warm till it 
 disappears. Where is the copper ? 4. Slide an iron nail into 
 
LABORATORY EXERCISES. 21 
 
 the tube, pour off the liquid at once, and rinse the nail ; what 
 is on it ? Which has the stronger attraction for nitric acid, 
 copper or iron ? What are some of the chemical properties 
 of copper? 
 
 Further Illustrations of Chemical Action. 
 
 6. Weigh the materials for i and 2. i. Heat i gram of 
 mercuric oxid in a clean, dry test tube. Make it very hot, 
 but move the tube about in the flame to avoid melting the 
 glass. Pass nearly to the bottom of the tube a splinter with 
 a spark upon the end. It will burst into flame ; if not at the 
 first trial, after removal and more heat. Now examine care- 
 fully the inside of the tube. Was mercuric oxid a simple 
 or compound substance } What is the nature of the two 
 products } This change illustrates decomposition. 2. Mix 
 thoroughly i gr. of zinc dust and J gr. of flowers of sulfur ; 
 (how do those weights stand related to the relative " atomic 
 weights " of zinc and sulfur i") place the mixture upon a piece 
 of iron (the ring stand) and apply a lighted match. Is the 
 result simple or compound .■• This change illustrates chemical 
 union or combination. 3. Weigh i gm. each of crystals 
 of barium chlorid and copper sulfate ; dissolve separately in 
 4C.C. of water, heating the water to boiling as the substances 
 dissolve. Now pour either into the other, drop by drop ; 
 -finally shake thoroughly; place the test tube in the rack. 
 After a few minutes notice that you have a clear green liquid 
 above and a heavy white powder below. The green liquid 
 contains copper chlorid and the white powder is barium 
 sulfate. The metals have changed places ; this kind of 
 change is called double decomposition. If you have at hand 
 a very good filter paper, you can separate these substances, 
 applying the knowledge of the process obtained in 3. 
 
Missing Page 
 
LABORATORY EXERCISES. 
 
 23 
 
 3. See whether you can get H from HCl and an iron nail. 
 (Warm.) 4. Arrange a side-necked test tube with cork and 
 delivery tube (Fig. 2), prepare H, using zinc, water, and 
 HCl ; let the H at first escape ; then collect over water 
 
 Fig. 2. 
 
 (teacher demonstrate) as follows : 5. Fill your larger beaker 
 and, lifting it mouth downward, pour up into the other 
 beaker, held mouth downward. Apply at once a light to 
 each beaker; where do you find the H ? Why.? 6. Fill a 
 test tube with H and light it. Fill a test tube two-thirds 
 with air, one-third with H, light it. Notice the difference 
 and remember never, under any circumstances, to apply a 
 light to a hydrogen generator or any of its connections until 
 you have satisfied yourself by collection of the gas in test 
 
24 
 
 EXERCISES IN CHEMISTRY. 
 
 tubes that pure hydrogen is coming away. 7. Fill your gas 
 jar with H ; lift from the tank and, holding it mouth down- 
 ward, light it with a burning stick. Will the stick burn if 
 
 carried into the jar above the 
 H flame ? Can you light the 
 stick at the H flame > What 
 gathers on the jar inside .' 
 What properties of H do 
 these experiments illustrate .'' 
 [8. Put in about 4 in. of i in. 
 glass tubing a little ferric 
 oxid, connect the tubing with 
 the H generator which must 
 be working well (Fig. 3). 
 Heat the tubing very hot 
 near the middle (see caution 
 in 6). As the H passes over 
 the hot ferric oxid it takes 
 away oxygen — reduces. Can 
 you discover any evidence of 
 this.?] 9. Evaporate some of the liquid from the side-necked 
 test tube. You find a white solid, zinc chlorid (ZnClg), which 
 will melt if you' heat too much. 
 
 Fig. 3. 
 
 Equations: (i) Zn+H2S04=ZnS04+ 
 H.+ . (3) Fe+2HC1= + 
 
 (8) rea03+6H=Fe2+ 
 
 (2&4; Zn+2HC1= 
 (6 & 7) H2+ =HsO. 
 
 0<cyg&n. 
 
 8. See 6, i. i. Heat in a dry test tube i gm. of potassic 
 chlorate (KCIO3). Test the gas evolved with a glowing 
 stick. Potassic chlorid (KCl) is left in the tube. 2. Dis- 
 solve a portion of this KCl in distilled water; dissolve in 
 
LABORATORY EXERCISES. 
 
 25 
 
 another portion of water a little unheated KCIO3; now to 
 each add a few drops of silver nitrate (AgNOg) solution. 
 This is merely to show that the salt after thorough heating 
 is different, and acts differently from the salt before heating. 
 3. Put in a (dry, clean) s. n. t. t. 3 or 4 gm. of KCIO3 and 
 about one-half as much manganese dioxid (MnOg), well 
 mixed. Arrange completely for the collection of a gas over 
 water, then heat the tube moderately. Remember that in 
 all such cases the first bubbles of gas are only the air in 
 the generator ; so let a little escape and then collect two 
 gas jars full. Immediately take the 
 delivery tube out of water in this and 
 all similar cases. (Why ?) Slip the 
 cover glasses on the jars before tak- 
 ing them out of water, in this and 
 all similar cases. (Why .') Take 
 them out and let them stand covered 
 till you are ready to use them (Fig. 4). 
 4. Into one jar (mouth upward or downward .' the 
 hydrogen density of O is 16, of air 14.44) plunge a 
 glowing stick ; when it bursts into flame remove it, 
 blow out the flame, return it, continue this till the O 
 is exhausted. 5. Wind a fine iron wire (No. 28) 
 around about |- of a match, using the phosphorus end 
 (Fig. 5). Now light the match and plunge it with the 
 wire nearly to the bottom of the other jar. A con- 
 siderable portion of the wire will burn brilliantly in the 
 jar of O. These two are typical of many cases in 
 which combustion or burning proceeds with much 
 '"''' '■ more energy in pure O than it does in air, which is 
 only ^ O. 6. Ozone (O3). Put about 2c.c. of dil. HgSO^ into 
 a test tube ; mix with it enough barium peroxid to make 
 
 Fig. 4. 
 
26 EXERCISES IN CHEMISTRY. 
 
 a thick paste ; finally warm. This reaction produces hydro- 
 gen peroxid (HgOj), oxygen, and a little ozone. Notice its 
 sharp, irritating odor, and the effect of the gas in the tube 
 (not the liquid) upon paper that has been dipped into a mix- 
 ture of iodid of potassium (KI) solution and starch paste. 
 (Starch and iodin form a blue compound. But starch and 
 KI do not unless something takes possession of the potas- 
 sium and releases the iodin. Ordinary oxygen cannot do 
 this, but ozone can. Hence this reaction is a test for ozone 
 when other substances that could take the potassium — like 
 chlorin — are absent.) 
 
 Equations: (l) KC10s=KCl-t- . (2) KCl+AgN03= 
 
 AgCl-f-KNOa. KC103+MnOa=Mn05+ + 
 
 Both equations represent only the final result. (5) 3Fe+? 0=Fe30i. 
 (6) Ba02+HjS04=Hj02+ mH20a=mHjO+nOa-|-p03 (values of 
 
 m, n & p unknown). 
 
 Water (JT,0). 
 
 9. See 6, 3, and 7, 7 & 8. i. Hold a dry cold beaker 
 inverted around the lamp flame an instant. 2. Heat a dozen 
 granules of sugar in a test tube ; does water appear ? [3. 
 Heat a little piece of meat in a test tube.] 4. Heat a crys- 
 tal of carbonate of soda ("washing soda") in a test tube. 
 What do I, 2, 3, 4 show, respectively, about water.' 5. Put 
 in a side-necked tube with clean delivery tube loc.c. of water, 
 boil with a small flame, and catch in a beaker a few c.c. of 
 the condensed steam. It is "distilled water"; its purity 
 depends on what was in the water at the outset. Why ? 
 You cannot get much this way. Why .■' See 28, i, and 
 Part IV, 6. Test the reaction (see 4) of distilled water. 6. 
 Test the reaction of water which has been shaken up with 
 the fumes of burning sulfur (ic.c. of water, one or two sulfur 
 
LABORATORY EXERCISES. 2/ 
 
 matches with the tip cut off). 7. Test the reaction of water 
 shaken up with a tiny lump of lime. 8. Test the reaction 
 of water (5c. c. or so) which has been boiled with xV S"^- 
 of lead oxid. 9. Test the reaction of water which has been 
 boiled with -jV gm. of arsenious oxid. (Arsenious oxid is 
 poisonous and must not be used carelessly. The reactions 
 in 8 & 9 are slow and feeble.) 10. Water of crystal- 
 lization. See 4. Heat in an evaporating dish a crystal of 
 copper sulfate, holding a dry cold beaker over it. What 
 appears on the beaker, and what do you hear and see in the 
 dish.' II. Efflorescence. Take a lump of fresh washing 
 soda, scrape off any white powder on its surface and let it 
 stand exposed to the air. In a dry room you will see some 
 change in ten minutes and a very great change in twenty- 
 four hours. It parts easily with its water of crystallization. 
 12. Deliquescence. Put on a piece of glass a little "fused" 
 (or anhydrous) calcic chlorid (CaCy ; see how the smallest 
 particles soon liquefy ; if, from the dryness of the air, this 
 does not take place in a few minutes, breathe on the CaClg. 
 For what can deliquescent substances be used .■' 
 
 Equations : (4) NajCOs, 10H2O=Na2CO3+ . (6) S+02= 
 
 S0.:+H20= (7) CaO+H20= (8) PbO+H20= 
 
 (9) AS2O3+? H,0=? HsAsOs. (10) CuSO,, 5H-,0=CuS04+ 
 
 Solution of Salts in Water. 
 
 10. In experiments i, 2, 3, use about i gm. of CUSO4, 
 5H2O, and about 8c. c. of water. In each case make two 
 solutions exactly alike, except in the condition in question. 
 I. Does copper sulfate dissolve more quickly with or without 
 agitation .' 2. Does it dissolve more quickly in hot or cold 
 water? 3. Does it dissolve more quickly powdered or in 
 
28 EXERCISES IN CHEMISTRY. 
 
 lump ? la 4 use carbonate of soda. 4. Does hot or cold 
 water dissolve more of this salt .■" Dissolve all possible in 
 8c.c. of cold water (cold saturated solution), now add nearly 
 as great a bulk of the carbonate as the bulk of the liquid ; 
 heat. 5. What happens as the solution (4) cools.' (If 
 nothing happens, it is because you failed to get a cold 
 saturated solution.) 6. Into 8c. c. of water put about i gm. 
 of plaster of Paris (CaSOJ ; you cannot see whether any 
 has dissolved or not. Filter a little of the liquid upon a 
 fragment of clean glass (not your cover glass) and gently 
 heat till the water evaporates. Had any of the CaSO^ 
 dissolved ? 7. A supersaturated solution. Fill a test tube 
 with small crystals of sodium thiosulfate (NagSjOg, 5H2O — 
 the photographer's "hypo"). Heat 2 or jc.c. of water 
 (about one-seventh the bulk of the " hypo " ), pour it upon 
 the crystals, dissolve them all by heating, bring to boiling, 
 cork securely, and allow the liquid to become cold again. 
 Then remove the cork ; if nothing happens, drop in one 
 very small crystal. Rapid crystallization will take place, 
 accompanied by a temperature change. 8. With ther- 
 mometer take the temperature of about 50c. c. of water in 
 a beaker, put in a spoonful of saltpeter (KNO3 potassium 
 nitrate) and take minimum temperature. 9. Into a beaker 
 put about soc.c. of water, boil upon the wire gauze and 
 note the temperature when fully boiling. Save for 10. 
 10. Now put in about two teaspoonfuls of common salt 
 (NaCl — sodium chlorid) and take the boiling temperature 
 of this solution. 
 
 Hydrogen Dioxid or Peroxid {HiOi). 
 
 11. Take about loc.c. to your table, i. Taste a few 
 drops. 2. Put about 2 gm. of MnO„ in a test tube and add 
 
LABORATORY EXERCISES. 29 
 
 5C.C. of HgOj ; prove that the escaping gas is oxygen. 
 3. To 2 or 3C.C. of solution of potassium permanganate 
 (KMnO^) add a few drops of dilute HgSO^ and then add 
 2 or 3C.C. of HgOj! prove that O escapes. 4. Gently warm 
 the remainder of the H^O^ with a little freshly made lead 
 sulfid (PbS). It gives O to the PbS, forming white PhSO^ 
 (lead sulfate). Does O escape ? 
 
 Equations: H20a=H20+0. (3) The action is complex; the final products 
 are K.SO,, MnSO,, H.,0, O. (4) ? H202+PbS=PbS04+ 
 
 iFluorin F) Hydrofluoric Acid {HF), 
 
 13. {Do not let the HF come in contact with the fingers.) 
 Coat a piece of glass about 2"xt," with paraffin ; scratch 
 through the paraffin some design upon the glass ; lay a piece 
 of filter paper upon the glass and wet it with HF ; after 
 fifteen minutes or more clean and examine the glass. To 
 clean the glass rinse with water, scrape with a knife, and 
 finally gently heat and wipe with a cloth. 
 
 Equations : Si02+? HF^? HjO+SiF4 represents in part the action. 
 
 * Chlorin {CI). 
 
 13. I. Make CI by MnOg, 5 or 6 gm. and concentrated 
 HCl, 5C.C., well mixed and slightly warmed in a test tube. 
 When the tube is half-full of CI (recognized by its greenish 
 color), heat the end of a narrow strip of copper foil and 
 plunge it quickly into the CI without letting the copper get 
 wet. If bubbles interfere, stop heating and wait for them 
 to subside. The copper glows and forms cupric chlorid. 
 Notice the color it produces in the lamp flame. Get rid of 
 
 *In paragraphs 13, 14, 15, the gases produced are irritating and hoods are desirable, 
 though with care it is possible to get along without them. 
 
30 EXERCISES IN CHEMISTRY. 
 
 the contents of the tube at once by rinsing into the drain- 
 pipe of the sink. 2. Make a less offensive solution of CI 
 by using about one teaspoonful of bleaching powder, soc.c. 
 of water, and adding sec. of dilute acid (acetic, hydrochloric, 
 or sulfuric). 3. Try the bleaching action of this solution 
 upon litmus paper, printer's and writing inks, and its effect 
 upon offensive material (HgS water may be used), and save 
 a little for 16, 4. 
 
 Equation: (i) ? HCl+Mn02=? H20+MnCl2+ , representing final 
 
 result. Cu-|-Cl2= . Bleaching and disinfecting are largely due to 
 
 these two reactions: H+C1= ? . HjO+Cl2=2HC!l+ 
 
 Hydrochloric Acid (MCI). 
 
 14. I. Put in a side-necked test tube, in the order given 
 (why ?), 4C.C. of water, 6c.c. concentrated HjSO^ and 3 or 4 
 gm. of coarse rock salt (NaCl) ; hold moistened litmus paper 
 at the end of the delivery tube. 2. Also breathe across the 
 end of the tube and notice the fog which is formed whenever 
 HCl gas comes in contact^with vapor of water. 3. After 
 the first rapid evolution of gas gently warm the s. n. tube, 
 if necessary (do not heat to the boiling point of water in 
 this or any similar case where a gas is generated in the 
 presence of water (why?)) ; now fill a perfectly dry gas jar 
 (mouth upwards or downwards i") with HCl gas. Remember 
 that you can get at the weight of a gas from its formula. 
 (See Part II, Sec. 9.) It is difficult to know when the jar is 
 full, so you must send a steady current of gas to the bottom 
 of the jar for at least 10 minutes, keeping the jar loosely 
 covered ; now cover securely, insert the jar in the water tank 
 and slip it off from the cover. 4. Moisten the inside of 
 your other gas jar with a few drops of strong NH^OH, and 
 
LABORATORY EXERCISES. 3 I 
 
 send HCl from the s. n. tube into this jar ; heavy clouds of 
 NH^Cl (a solid). Why not collect HCl over water as you 
 did O and H .? 
 
 Equations: (i) NaCl+H2S04=HNaS04+ . 2NaCl+ 
 
 HiiS04=Na3S04+ . (3) NHa+HCl= 
 
 Bromin (,Br). 
 
 15. Upon I crystal of potassium bromid (KBr) put 2 or 
 3 drops of concentrated HgSO^, and gently warm ; reddish 
 brown Br and colorless fuming HBr are evolved. 2. Nearly 
 fill a test tube with water ; add i drop of Br (this must be 
 done with a dropping tube and in a hood), shake the tube 
 to show the solubility of the Br. 3. Try the bleaching 
 power of a little of this Br water upon litmus. 4. Transfer 
 the Br water to a small beaker and pass HgS (hydrogen 
 sulfid) rapidly through it. The HgS is made in the s. n. test 
 tube by the action of quite dilute H2S0^ upon 2 or 3 gm. 
 of FeS (ferrous sulfid). Notice the change in color and 
 clearness and the effect upon litmus (compare with 3). HBr 
 is formed and S is set free. 
 
 Equations: (i) like 14, 1, but 2HBr+H,S04=? H,0+Br,+SOs. 
 
 (4) perhaps 10Br+2H,S+4HaO=10HBr+H2SO4+S. 
 
 lodin (J). 
 
 16. I. Upon one crystal of KI put 2 or 3 drops of con- 
 centrated H2SO4 and gently warm ; I and HI are evolved. 
 What is the color of I vapor ? 2. Put one small scale of I 
 into each of four test tubes ; then put into the tubes respect- 
 ively 2C.C. of water, alcohol, carbon disulfid (CS^, water 
 solution of KI. In which of the liquids does the I dissolve, 
 and what are the colors .' 3. To a beaker of dilute starch 
 
32 EXERCISES IN CHEMISTRY. 
 
 paste (see Chemicals) add a drop or two of either the 
 alcoholic or the KI solution of I ; here we have free I. 
 What is the effect on the starch ? 4. Add to another lot of 
 starch paste a drop or two of fresh KI solution. In this 
 there should be no free I. What is the effect on the starch } 
 Now add CI water from 13, 3, first a drop, then loc.c. or 
 more ; the liberated I forms a blue compound with starch, 
 which more CI destroys. 5. Test a sample of baking powder 
 for starch. 
 
 Equations: (l) like 15, 1. (4) KI+C1= + 
 
 Sulfur (S). 
 
 11. I. Put a piece of broken roll sulfur, as large as half 
 a pea, into 2c.c. of CSj and shake for a few minutes. Does 
 it perceptibly dissolve .-' Could a substance be soluble and 
 yet not wholly dissolve ? Why ? 2. Pour the liquid (from i) 
 upon the dry bottom of an inverted gas jar, and let it 
 evaporate spontaneously. Notice shape of crystals. 3. 
 Nearly fill a 2 oz. hessian crucible with broken roll sulfur 
 and heat until all has melted. Do not allow it to burn at 
 the top or sides of the crucible. Let it cool uncovered. 
 As soon as the surface is just crusted over (watch it care- 
 fully), make a hole and quickly drain out all the liquid. Cut 
 out the crust. Are the crystals like those in 2 1 4. Put 
 some fragments of roll sulfur into a large test tube (not side- 
 necked), and apply heat, noting carefully changes in color 
 and thickness which take place as the temperature rises. 
 5. When the sulfur is boiling, pour part of it into a beaker 
 of cold water, and examine. Recall 5, 2 and 6, 2. 6. 
 Ignite a little S. Does it take fire easily.? Describe the 
 flame. Where is the product of the combustion .? Can you 
 smell it } 7. Hold a strip of moistened blue litmus paper in 
 
LABORATORY EXERCISES. 33 
 
 the fumes of the burning S. 8. Make a simple blow-pipe by 
 bending about 8 in. of -^g glass tubing at a right angle i in. 
 from the end and closing the shorter arm to about the 
 caliber of a pin. Upon a piece of charcoal put a mixture of 
 one part finely powdered FeSand four parts of dry Na2C03; 
 the whole mixture should not measure more than ice. ; 
 fuse thoroughly with the blow-pipe flame. Put the fused mass 
 upon a silver coin and moisten with water. Sodium sulfid is 
 formed by the fusion, and this with the silver makes black 
 silver sulfid. Give the history of the sulfur in this experiment. 
 Equations: (6) S+02= . (7) S02+H20= 
 
 Hydrogen Sulfid (H^S). 
 
 18. Have everything ready for the whole paragraph before 
 you generate HjS. i. Prepare HgS as in 15, 4. 2. Notice 
 odor. 3. Direct the gas upon moistened litmus paper, red 
 and blue, until you notice some change. 4. Direct the gas 
 upon paper which has been wet in solution of acetate or 
 nitrate of lead ; the black product is PbS. 5. Collect a little 
 H2S over water — a test tube full — apply a light. It burns, 
 with almost invisible flame, depositing some sulfur. 6. Mix 
 in a test tube solutions of antimony chlorid (SbClg) and 
 ferrous chlorid (FeCy. Make them distinctly acid with 
 dilute HCl, if they are not already, and pass HgS through; 
 orange antimony sulfid (SbjSg) falls as an insoluble com- 
 pound, or precipitate, but no FeS falls. Now filter and add 
 NH^OH to the filtrate till alkaline after stirring; in the 
 presence of the NH4OH black FeS is formed and pre- 
 cipitated. All the metals that act like antimony can be 
 separated from all the metals that act like iron in this way; 
 and this and similar methods are largely used in the appli- 
 cation of chemistry, called Qualitative Analysis. 
 
34 EXERCISES IN CHEMISTRY. 
 
 Equations: (l) FeS+H,SO»= + . (4) H,S+Pb(N03)5= 
 
 ?HN03+ . (5) H,S+? 0=H20+S02 and H2S+0=S+ 
 
 (6) ? SbCl3+? H,S=? Sb.S3+? HCl. NH40H+HC1=NH4CH- 
 ? NH40H+H2S=(NH4)2S+ . (NH4)i,S+FeCl2= + 
 
 Sulfurous Oxid (SO2). 
 
 19. See 17, 6. Before generating SOg have all the 
 materials for the paragraph ready, i. Put in a side-necked 
 test tube i sq. in. of sheet copper, not more than locc. of 
 concentrated H2SO4 and apply heat until the gas begins to 
 come, then regulate the heat carefully. On no account dip 
 the delivery tube into any liquid, as the sucking back of 
 water upon the hot HgSO^ would produce a dangerous ex- 
 plosion. Get information as to odor, color, relative weight, 
 reaction, and combustibility of SO2. 2. Put in a gas jar a 
 few c.c. of logwood solution, let the SO2 run into the jar 
 (mouth upwards), then cover and shake. [3. If obtainable, 
 a blue or purple flower may be put in the jar instead of 
 logwood to show bleaching effect.] 4. Put in another jar a 
 dilute solution of KMnO^ and let the SOg run into this jar 
 as in 2. The SOj takes oxygen, finally forming sulfates of 
 potassium and manganese. Save part of the liquid in the 
 jar for 30, 7. 5. Let the contents of the s. n. test tube 
 cool, pour out the liquid and remove any copper that is left ; 
 what do you find at the bottorn of the tube ? Why is it not 
 blue.? What is the effect on the color of adding water? 
 6. What gas is evolved by the action of dilute HjSO^ upon 
 a little solid sulfite of sodium (NagSOg) in a test tube ? 
 
 Equations; (i) final result Cu-|-? HjS04=CuS04-f- -f-? HjO. 
 
 (4) perhaps 3S02-|-2KMn04=K2S04-F2MnS04H-02. (6) NasSOa-h 
 HsS04=:Na2S04-|- + 
 
LABORATORY EXERCISES. 35 
 
 Sulfuric Acid {H^SOt). 
 
 30. I. Put about 2C.C. of water in a test tube, and add 
 an equal volume of concentrated HjSO^, holding the lower 
 part of the tube in the hand. (Remember in diluting con- 
 centrated H2SO4 always to add the acid to the water, and 
 slowly.) 2. Make a very thick syrup of sugar and hot water 
 (test tube), and after it has nearly cooled add >n equal vol- 
 ume of concentrated H2SO4, holding the tube by a test tube 
 holder. 3. Trace some design upon white paper with dilute 
 H2SO4, and gently warm the paper until it is dry. 4. Add 
 dilute HCl and then baric chlorid solution (BaClg) to dilute 
 H2SO4. 5. Repeat 4, rising a solution of NagSO^ instead of 
 H2SO4. The precipitate in 4 and 5 is barium sulfate (BaSO^), 
 and this is a common method of identifying sulfates. The 
 HCl has no effect here, but without the HCl a number of 
 other substances — carbonates, phosphates etc. — might be 
 mistaken for sulfates. 6. Try phosphate of soda solution 
 and BaCl2, with and without HCl. 7. Test the liquid saved 
 from 19, 4, for sulfates. [8. Repeat 15, 4, and ascertain 
 whether any sulfuric acid is formed.] Very few illustrations 
 of the properties of this important substance are given here 
 because illustrations are all through the book. How do 
 experiments 7, i, 8, 6, 14, i, 15, i, 16, i, 18, i, 19, i, 4, 6, 
 illustrate properties of H2SO4.'' After you have done the 
 following experiments, note here the properties of HgSO^ 
 which they illustrate: 33, 5, 38, i, 2, 36, 6, 38, i, 
 39, I, 47. 
 
 Equations: (l) perhaps H2SO4 + 2H20 = H8S06. (2) Ci^H^sOi, is 
 
 partly broken into carbon and water. (4) H2SO4 + BaCla + nHCl = nHCl 
 + BaSO, + • (5) Na,S04 + BaCl, + nHCl = 
 
 + + 
 
36 
 
 EXERCISES IN CHEMISTRY. 
 
 Nitrogen (N), 
 2i . Put in a s. n. test tube 2 gm. of NH^Cl, 3 gm. of 
 potassium nitrite (KNOj), just moisten with a few drops of 
 water. Apply gentle heat and collect a jar of N over water. 
 Does N maintain combustion or burn ? 
 
 Equations: NH4CI + KNO^ = NH4NO, + • NHjNOi 
 
 Air, 
 
 33. Take about 8 in. of one-fourth or five-eighths in. glass 
 tubing; close one end and bend about one inch to a right 
 
 angle ; bend the 
 open end in a 
 short loop (Fig. 6). 
 After the tube is 
 cool slip to the 
 closed end a piece 
 of phosphorus (P) 
 as large as a grain 
 of rice. (Dry the 
 P with filter paper, 
 do not rub or 
 squeeze it, and do 
 not at any time 
 touch it with the 
 fingers. Why .■' ) 
 Fill the open loop 
 with water, mark 
 its height with a file, then stand the tube upright in a 
 beaker of water, the open end half an inch or so under 
 water. Now heat the P gently for several minutes, and 
 
 Fig. 6. 
 
LABORATORY EXERCISES. 37 
 
 let the whole stand until it is cold again. (Why .?) Observe 
 that water has filled almost exactly one-fifth of the 
 air space in the tube. The P has exhausted the oxygen. 
 
 2. Demonstrate the presence of carbon dioxid (COj) in the 
 air by exposing half an hour Or more loc.c. of lime water 
 (Ca(0H)2) in an open beaker ; CaCOj is formed on the surface. 
 
 3. Half-fill two test tubes with beef tea. Heat both to boil- 
 ing, then plug the mouth of one loosely with a wad of cotton 
 that has just been passed through the lamp flame two or 
 three times. Germs from the atmosphere will cause putre- 
 faction and mould in the open tube but not in the plugged 
 tube. In the laboratory this process is likely to be very 
 slow — two or three weeks. 
 
 Equations: (i) P.; + ? = P206. P205-|-H20 = ? HPOs, or, 
 
 PsOs-)-? H,0 = 2HsP0.. (2) CaO-,H2-|-COj= -f 
 
 Ammonia (iVHs). 
 33. I. Grind together in a mortar about 2 gm. of NH^Cl 
 and 4 gm. of CaO ; NH3 is set free, to be recognized by its 
 odor, its effect upon moist red litmus paper and the fumes 
 when a glass rod moistened with concentrated HCl is brought 
 near. Save for 3. 2. Put in a test tube a solution of any 
 ammonium salt (NH^Cl or {NH^)2S04), add a few drops of 
 NaOH solution and warm ; is NHg liberated ? How can 
 you tell NH3 from any other gas ? 3. Transfer the mixture 
 from I to a s. n. test tube ; just moisten it and collect, with 
 heat, a dry jar of NHg and show its absorption by water. 
 The experiment is exactly like the last part of 14, 3, except, 
 which way should the mouth of the jar be — up or down — 
 to collect a gas like NHg ? [4. Get a s. n. test tube all ready 
 for the preparation of O (8, 3) ; put in a beaker about 20c.c. 
 
38 
 
 EXERCISES IN CHEMISTRY. 
 
 of strong ammonium hydroxid (NH^OH) ; set it upon the 
 sheet iron or gauze and gently heat till bubbles of ammo- 
 nia come rapidly 
 away, then make 
 O not too fast 
 and have it pass 
 through the liquid 
 in the beaker ; 
 when both gases 
 are coming well, 
 apply -a light to 
 the mouth of the 
 beaker. If you 
 have managed 
 this well, you will 
 have a brilliant 
 combustion of O 
 & NH3.] 5. Neu- 
 tralization. Re- 
 call 14, 4. To 
 loc.c. of dilute 
 HjSO^ add dilute 
 NH^OH until neither red nor blue litmus is affected by the 
 liquid. Use only one piece of litmus paper until the end of 
 the operation. Try not to use too much NH^OH ; but if you 
 do, a few drops of the acid may be added, etc. Save for 
 34, 2. 
 
 Equations: (i) CaO -|- ? 
 HC1 + NH3 = 
 Na.,S04 + ? H,0, 
 
 NHjCl = 2NH3 + CaCl^ + 
 (2) (NH02SO4+? NaOH = ? NHa + 
 NH^Cl-fNaOH= + (4)?NH,+ 
 
 = ? H,0-|-? N. (5) 2NH,OH4-H,SO4=(NH02SO,+ 
 
LABORATORY EXERCISES. 39 
 
 Compounds of Atnmonittm (IfBi). 
 
 34. I. To ICC. of solutions of chlorid of cadmium and 
 chlorid of tin, in separate test tubes, add first 3 or 4 drops, 
 and' then loc.c. of ammonium sulfid ((NH^)2S) ; warm ; in 
 both cases a sulfid is formed, but the tin sulfid dissolves in 
 the excess and the cadmium sulfid does not. A similar reac- 
 tion is used in Qualitative analysis to separate several 
 metals that act like tin from several that act like cadmium. 
 2. Evaporate the neutral solution of (NH4)2S04 (33, 5) 
 nearly to dryness, let it cool ; small crystals of (NH4)2S04 
 will form. See Part IV, 9. In a manner similar to this, 
 most ammonium salts can be made. How could you make 
 NH^NOg .' Examine the various ammonium salts at hand ; 
 do they smell of NH3 ? 
 
 NUrogen Monoxid {W^O). 
 
 35. Collect over water one jar of N^O, made by heating 
 about 4 gm. of NH^NOg in a s. n. test tube. The NH^NOg 
 melts to a liquid before the gas comes ; after the evolution of 
 gas begins, use just enough heat to keep up a steady current. 
 Refer again to cautions in 8, 3. Observe sweetish taste and 
 odor and effect upon combustion (stick). 
 
 Equations: NH4N03 = + . Compare with the decomposition 
 
 of NHiNOs in 21. In combustion N2O = Nj + O. 
 
 Nitrogen Dioacid {NO). 
 
 36. I. Equip a 250C.C. Erlenmeyer flask with delivery 
 tube and funnel tube, as in Fig. 8. See that all connections 
 are tight. Slide into the flask (do not drop it in) about 2 
 sq. in. (6 or 7 gm.) of sheet copper ; now add through the 
 funnel tube 20-30 c.c. of HNO3 sp. gr. 1.2. The action may 
 
40 
 
 EXERCISES IN CHEMISTRY. 
 
 be started by warming, but do not continue the heat. Collect 
 two jars of the gas ; one is to be used in 2 and must be full ; 
 the other is to be used in 3'? and need not be full, and is not 
 
 Fig. 8. 
 
 to be taken from the water tank. 2. Remove the full jar 
 from the tank ; slide the cover off the least possible degree 
 and pour in i c.c. (and not a drop more) of CSg. Shake 
 thoroughly, uncover completely, and at the same instant 
 apply a light. (This " flashlight " affects a photographic 
 plate in a dark room.) 
 
 Equations: (l) 3Cu + 8HN03=? Cu(NOs)2 + ? NO + ? H^O. 
 (2) Something like this : CS2 + 4NO = 00^ +30, +4N + S. 
 
 Kitrogen Tetroncid (JVO2J. 
 
 37. Arrange for the preparation of O (8, 3). Send 
 the O very slowly and intermittently into the jar of NO 
 from 36. Notice the immediate formation of the brown 
 gas (NO2) and its rapid absorption by water. If you manage 
 
LABORATORY EXERCISES. 
 
 41 
 
 this skilfully, you can carry the water nearly to the top of 
 the jar. 
 
 Equation: NO + 0=NOj, 
 
 Nitric Acid {HNOz). 
 
 38. [i. Where the conditions are suitable a little strong 
 HNO3 may be made by putting 20 gm. of common saltpeter 
 (KNO3) and 20 gm. of concentrated H2SO4 into a 2 to 4 oz. 
 glass-stoppered retort (Fig. 
 9) and gently distilling in 
 the hood. This must not be 
 done in the open laboratory 
 or with corks and rubber 
 connections.] 2. Into a test 
 tube put 2C.C. solution of 
 KNO3, ICC. concentrated 
 HjSO^, and a strip of copper 
 foil. This makes and uses 
 up HNO3 in the same oper- 
 ation. Nitrates may gener- 
 ally be recognized by this 
 reaction. 3. liito 2c.c. of concentrated HNO3 put a scrap 
 of tin ; the tin is oxidized to a white, nearly insoluble powder 
 (HjSnOg, metastannic acid), while in 2 and 36, i the 
 copper enters into the formation of a soluble nitrate. 4. 
 Boil a little finely powdered FeS in about 2c.c. of concen- 
 trated HNO3 carefully for two or three minutes, dilute with 
 water to 10 or 15C.C., filter, and add solution of BaClj or 
 . Ba(N03)2. A white precipitate. What does it signify } 
 See 30, 4 and 5. [4. Put in test tube about 4c. c. concen- 
 trated HNO3 ; loosely plug the mouth of the tube with horse- 
 
 Fig. 9. 
 
42 EXERCISES IN CHEMISTRY. 
 
 hair, bring the HNO3 to boiling till its vapor comes in con- 
 tact with the horsehair. Hood only.J 
 
 Equations: (i) 2KNOs+HsS04 = ? HNO3 + • With less 
 
 heat and nitrate, KNO3 + H2SO4 = HKSO4 + . (z) like I and 
 
 26, 1. 
 
 Phosphorus (1*). 
 
 39. (See cautions in 33, i. Do not let it become dry 
 until you wish to ignite it. Have water at hand to extin- 
 guish it if it inflames prematurely. It can safely be cut or 
 handled under water), i. Take a piece of yellow P the 
 size of a grain of rice, partially dry it ; notice fumes and 
 odor. 2. Put it into 4 or 5c.c. of water and warm. Notice 
 its low melting point — about 44°c. 3. Pour it into cold 
 water, cut it in two, dry half, dissolve it in 2c.c. of CSg (dry 
 test tube). This is its best solvent. 4. Pour this solution 
 upon a strip of filter paper, letting the drops fall into the 
 sink and leaving enough of the paper dry to hold it by; 
 wave the paper back and forth ; the CSj evaporates and the 
 finely divided P ignites spontaneously. Notice that generally 
 the paper does not take fire, which illustrates the low 
 kindling point of P. 5. Dry the other half of the P, put 
 it on a crucible cover, ignite it with a warm glass rod and 
 cover at once with a perfectly dry beaker. The success of 
 this depends upon the dryness of everything. See the 
 white, flaky product — P205- 6. Notice also how deliques- 
 cent this oxid is ; breathe on it. 7. Wash the inside of the 
 beaker with a few c.c. of water. Is this solution strongly 
 acid to litmus 1 It contains HPO3 metaphosphoric acid. 
 [8. Does " red " or " amorphous " P melt, dissolve, or ignite 
 as the other form did.' ] 
 
 Equations: 2P-|-? 0=Pa05. PjOs -f HjO = ? HPO3. 
 
LABORATORY EXERCISES. 43 
 
 Hydrogen Phosphid (Jffa-P). 
 
 [30. This gas may be made by a number of methods, of 
 which the simplest is : Drop into dilute HCl in a beaker or 
 conical test glass pieces of calcium phosphid (Ca3P2) of the 
 size of a pea until it comes well. Notice odor and spontane- 
 ous inflammability. (Hood only.) Let the pupil find other 
 methods in the text-books. 
 
 Equation: Ca^P, + 6HC1 = ? CaCl, -f ."] 
 
 Hypophosphorous Acid (ZTsPOa). 
 
 31. I. To 3C.C. of solution of CuSO^ add 2c.c. dilute 
 HgPOg and gently warm until no more change occurs. 
 During the change are formed cuprous hydrid (CU2H2), 
 cupric hydrid (CuHj.), copper (Cu), phosphorous acid 
 (H3PO3). It is a powerful reducing agent. See 7, 8. 
 
 2. Try in the same manner the reducing action of H3PO2 
 upon AgNOg solution. AgH is produced which, upon char- 
 coal with the blow-pipe, will give a little globule of silver. 
 
 The Phosphoric Acids (^HPO,, HiP^O-,, H^POi). 
 
 [33. I. Put about 2 gm. of "glacial" or metaphosphoric 
 acid (HPO3) into an evaporating dish, add 2 or 3c.c. of water, 
 and after a little of the HPO3 has dissolved pour the liquid 
 (only) into a test tube, now add AgNOg solution ; white 
 precipitate of AgPOg. 2. Now to the HPO3 left in the dish 
 add 20 to 30C.C. of water, dissolve with heat and boil a 
 minute or two after complete solution. Take a little of the 
 liquid out and add AgNOg ; little or no precipitate until you 
 add a drop or two of NH^OH, when white Ag^P^O^ appears. 
 
 3. Now fill the dish with water, keep it full, and boil half an 
 hour or more : test a portion of this solution with AgNOg, 
 
44 EXERCISES IN CHEMISTRY. 
 
 adding the merest drop of NH^OH and you will get light 
 yellow AggPO^. These experiments mark the progress 
 from HPO3 to HgPO^ by combination with water. All pi 
 the precipitates are soluble in NH^OH, so you must be 
 careful not to add much. 
 
 Equations: ? HPO3 + H.O = H,P,0,. H4P50,+H20=? HsPO., 
 
 HPO3+ =H3PO.. HP03 + AgN03= + H1P2O, 
 
 + ? AgNO, = + . H3PO,, + ? AgNOa = 
 
 + •] 
 
 Arsenic (^s). 
 
 33. Examine the element arsenic and such compounds 
 as are on the shelf, i. Draw out 8 in. of i in. glass tubing 
 so as to make two 4 in. tubes each closed at one end. 
 Convey to the closed end of one of the tubes about ^ gm. 
 of white arsenic (arsenious oxid, AsjOg) — about what you 
 could put on the tip of a penknife. Now heat and notice 
 that the AsgOg becomes a vapor and condenses in glittering 
 (crystalline) form on the cooler portion of the tube (it 
 sublimes). Do not heat long enough to drive any vapor out 
 of the tube, as it is very poisonous. Save tube for 3. 2. 
 Intimately mix the same amount of AsgOg with powdered 
 charcoal, put it in the other tube, put about an inch of 
 charcoal above it, then heat first the charcoal, then the 
 mixture. You will need the heat of a Bunsen burner here. 
 After sufficient heat you will find on the tube a dark shining 
 band or "mirror" of elemental arsenic, reduced by the 
 charcoal. 3. Cut off the portion of the tube (i) containing 
 the AsgOg and boil it in 10 to 20c. c. of water in the evap- 
 orating dish ; test its reaction (see 9, 9). 4. Now add a 
 few drops of HCl and a slip of bright copper foil and boil 
 again ; the arsenic is deposited on the copper, alloying 
 
LABORATORY EXERCISES. 45 
 
 with it. 5. Dry the copper thoroughly with filter paper and 
 heat it in a perfectly clean and bright test tube ; a faint 
 ring of AsgOj crystals appears just above the hot portion of 
 the tube. (See what you can find about the Marsh-Berze- 
 lius test for arsenic). 
 
 Equations : (2) As^Os + ? C = ? CO + 2As. (3) As^Oa + 
 
 ? H,0 = ? HsAsOs. ' (4) H3Aa03+? HCl=AsCl, + 
 2AsCls + 3Cu = ? CuCla + 2As. ? As + ? O == AssOa. 
 
 Antimony (Sb). 
 
 34. I. Upon a piece of Sb as large as a grain of rice try 
 the efEect of hot concentrated HCl. 2. After rinsing try 
 upon the same piece the effect of concentrated HNO3. Prod- 
 uct is HSbOg. Compare 38, 3. 3. Dissolve with heat i 
 gm. of SbgSg in 3 or 4C.C. concentrated HCl ; what gas 
 escapes ? Evaporate tlie solution nearly to dryness ; the 
 product is SbClg. 4. Pour this thick liquid into a beaker 
 of water; white SbOCl is produced. Occasionally, when 
 the evaporation is insufficient, an orange instead of a white 
 precipitate is obtained in 4. What is it ? See 18, 6. 
 
 Equations: (3) Sb,S, + ? HCl = ? SbCl3+ . (4) SbCla 
 
 + H20 = SbOCl + 
 
 Boric Acid (XTaBO,). 
 
 35. I. Make a loop in a platinum wire by winding one 
 end 01 it halfway up the tip of a lead-pencil. Melt into this 
 loop little bits of borax until the loop is just filled with a 
 clear glassy substance. Now melt thoroughly into this glass 
 with your blow-pipe, pieces no larger than the head of a pin, 
 of salts of either iron, copper, cobalt, nickel, manganese, 
 chromium. Each of these gives a characteristic color. If 
 
46 EXERCISES IN CHEMISTRY. 
 
 you have not the platinum wires, you may use an iron wire, 
 but the bead will then be colored at the start from the iron. 
 Look up "Borax Bead." 2. Using about loc.c. of water, 
 make a hot saturated (?) solution of borax (Na^B^O^ is " dry " 
 borax) in evaporating dish ; transfer it to a beaker and add 
 concentrated HCl hi excess (.'). H3BO3 separates in pearly 
 scalelike crystals ; when the mixture is cold transfer it to a 
 filter and wash three or four times with a very little cold 
 water. 3. Try the solubility of some of these crystals in 
 boiling water. 4. Put loc.c. of alcohol in an evaporating 
 dish ; add some of the H3BO3 crystals, ignite the alcohol, 
 stir and notice the green flame characteristic of H3BO3. 
 (The flame may be tinged with yellow from the presence of 
 the sodium.) 
 
 Equation: (2) Na^BiO, + ? HCl +? H^O = ? NaCl + ? H,B03. 
 
 Compounds of Silicon (Si). 
 
 36. I. To ICC. of solution sodium silicate (Na^SiOJ add 
 an equal amount of concentrated HCl; gelatinous H^SiO^ 
 (silicic acid) forms at once. 2. Put in a beaker about i6c.c. 
 of water, add just half as much Na^SiO^ solution, and acidu- 
 late with dilute HCl. H4Si04 remains in solution. (What 
 is the other product i") 3. Transfer the mixture (2) to an 
 evaporating dish and bring to dryness. At a certain stage 
 the liquid will become gelatinous ; at this point the evapora- 
 tion must be carried on with constant stirring and very gentle 
 heat, or everything will jump out of the dish. The end is 
 reached when a dry white powder is left in the dish. After 
 the dish has cooled, add water, filter, wash the precipitate 
 and dry it in the paper. Gritty, white Si02 is obtained. 
 What ran through in the filtrate .'' Prove by tasting. 4. 
 Try the solubility of this SiOg in boiling NaOH solution. 
 
LABORATORY EXERCISKS. 47 
 
 5. Acidulate liquid from 4 with HCI. It contains Na4Si04 
 and will act like i or 2 according to its strength. Why do 
 the stoppers of NaOH bottles stick ? [6. Take pure sand, 
 or powdered glass, or SiOg made as in 2 & 3, about 2 gm., 
 fluor-spar 2 gm., mix, put in a test tube, add enough concen- 
 trated H2SO4 to make a paste, adjust cork and delivery tube ; 
 put in another test tube ^ in. of mercury ; dip the delivery 
 tube iilto the mercury and keep it there; nearly fill the 
 second tube with water, finally gently heat the first tube. 
 SiF^ is evolved ; as it passes into the water it forms hydro- 
 fluor-silicic acid (HjSiFg) and gelatinous H^SiO^. Save the 
 mercury. 7. Filter some of the liquid (6), and notice that 
 the filtrate is strongly acid. Add to it a solution of some 
 potassium salt, e. g. KCl ; precipitate of KgSiFg (not in very 
 dilute solutions).] 
 
 Equations: (i) NaiSiOi + ? HCI = H.SiO, + (3) HiSiO, 
 
 — H,0=: . HiSiO, — 2H,0= . (4) SiOa + ? NaOH 
 
 = Na4Si04+ . (6) CaT^ + HaSO, = + SiO^ 
 
 + ? HP= +? H2O. ? SiF4 + ? H,0 = H4Si04 + 
 
 ? HsSiPe. (7) H^SiFa + ? KCl = ? HCI + 
 
 Carbon (C). 
 
 37. Examine the various forms of carbon, i. Put about 
 2 teaspoonfuls of powdered bone black into a filter, making 
 the center concave ; filter slowly through this 20 or 30C.C. of 
 logwood solution. 2. Upon a piece of charcoal put ^ gm. or 
 less of PbO and direct with the blow-pipe the inner lamp 
 flame upon it. Illustrates reducing action of C. See 33, 2 ; 
 also compare with 7, 8, and 31, i, 2. 
 
 Equation: (2) PbO + C = CO+Pb. CO + O = 
 
48 EXERCISES IN CHEMISTRY. 
 
 Carbon Monoxid iCO). 
 
 38. I. Put in a s. n. test tube 2c.c. of formic acid (HgCOj) 
 and 4 or 5c.c. of concentrated H2SO4; gently warm. Collect 
 one jar of the gas over water for 2. Pass the gas through 
 lime water — about 20c.c. in a beaker ; no result. 2. Remov- 
 ing the jar (i) covered, uncover and apply a light. Where 
 have you seen this blue flame before 1 3. Pour the lime 
 water (i) into the jar and shake ; it turns milky now, show- 
 ing the presence of COg. 
 
 Equations: (l) H^COj + HjSO, = (H2SO4 -f H^O) -f CO. (2) CO -|- 
 
 O = . For 3 see next paragraph. 
 
 Carbon Dioxid (COa). 
 
 39. I. Put a small crystal of NagCOg, loHjO into each 
 of two test tubes ; to one add dilute HCl, to the other dilute 
 H2SO4 ; rapid escape of CO^. What are the other products ? 
 2. Put in the Erlenmeyer flask (36, i) about lo gm. of 
 marble dust (CaCOg) ; cover with water and add concentrated 
 HCl (lo to 20 c.c), until the gas comes well. Put about 
 20c. c. of lime water (Ca(0H)2) in a beaker, test its reaction, 
 and make the COg pass through it until it becomes first 
 milky and finally clear or almost clear again. Try the reac- 
 tion when it is clear. There is produced first insoluble 
 CaCOg and later soluble H2Ca(C03)2. Save for 3. Collect 
 two jars of CO2 over water. Remove one, covered, turn the 
 mouth the right way (.■'), uncover and plunge into the jar a 
 splinter which has been dipped in turpentine and lighted. 
 4. Hold a lighted match near the surface of the table and 
 pour the gas in the other jar rapidly upon it. 5. Put the 
 beaker (2) on the gauze and boil the liquid. As the excess of 
 CO2 is boiled out, CaCOg is again deposited, though in a 
 
LABORATORY EXERCISES. 
 
 49 
 
 I. In iG.c. of CSg shake for a minute or two 
 
 denser form than in 2. 6. Hold for an instant a beaker in- 
 verted above the lamp flame, then test the air in it for COj 
 (with a little Ca(0H)2). 7. Lower a little piece of burning 
 wood into a gas jar and test the air in it for COg. 8. Prove 
 that the air that comes from the lungs contains COj. 
 
 Equations: (i) NajCOa + HCl = HNaCOa + , in very dilute solu- 
 
 tions. NajCOa + 2HC1 = H2O + -f- NajCOs+H^SO, 
 
 = NaaS04 + + . (2) CaO,H2 + CO^ = CaCOs + 
 
 CaCOa + COi, -f =HjCa(C03')i,. (5) HjCaCCOO^ = CaCOa -f 
 
 H,0-f 
 
 Carbon Disulfid iCS%)m 
 
 40. Recall its use as a solvent in 16, 2, 17, i, 39, 3 ; 
 also see 36, 2, 
 
 a little chip of caoutchouc ; pour 
 the liquid upon the bottom of 
 an inverted gas jar or beaker 
 and let it evaporate. A film of 
 rubber is left. 2. Into a dry jar 
 put ICC (not anymore) of CSg ; 
 cover, shake, and plunge into the 
 jar a hot glass rod. This shows 
 the low kindling point of CSg. 
 What gas can you smell in the 
 jar 1 [3. Take about 6 in. of 
 J in. glass tubing, close one end, 
 and bend near the closed end 
 like Fig. 10. When it is cool 
 pour to the closed end ic.c of 
 
 i 
 
 CSj. Keep hqrizontal and con- 
 
 Fig. 10. 
 
 vey about i gm. of zinc dust to 
 
 the middle of the tube. Now heat the Zn very hot. Keep it 
 
 hot, then cautiously wave the flame under the CSg till its 
 
50 EXERCISES IN CHEMISTRY. 
 
 vapor passes over the Zn. Do not put much heat on the CSg 
 or it will suddenly jump out of the tube. ZnS is formed 
 with a glow and C deposits on the tube. Can you get HgS 
 from the ZnS and HCl ? See 18, 4.] 
 
 Equations: (2) CS, + ? O = COs + SO, + S, perhaps. (3) CS, + 
 2Zn = ? ZnS + . ZnS + 2HC1 = + 
 
 The Lamp Flame. 
 
 41. I. Examine the lamp or candle flame and identify 
 three somewhat distinct cones. 2. Holding a piece of white 
 paper horizontally, carry it down upon the flame nearly to 
 the wick and bring it up and away before it can take fire. 
 The paper is charred in a circle. 3. Hold a thin stick across 
 the flame and remove it before it burns. It is charred in two 
 places, the part within the flame being affected not at all, 
 or very little. 4. Quickly carry the phosphorus end of a 
 match through the flame to the center of the wick, and with 
 great quickness remove it. With skill enough the match 
 can be brought out without igniting. $. Take 2 in. of ^ in. 
 glass tubing, hold it by tongs with the open end in the 
 center of the flame at an angle of about 45"; see if you can 
 light the vapor that comes through the tube. From these 
 experiments what do you learn about the lamp flame ? 6. 
 Examine the Bunsen burner, and try it, getting the luminous 
 smoky flame and the blue non-luminous flame. 7. While 
 you have the Bunsen burner (luminous flame), carry some 
 wire gauze halfway down the flame ; does the flame pass 
 through .'' 8. Try blowing out the gas and fighting it be- 
 tween two pieces of gauze held horizontally. 9. Do 7 and 8 
 work as well with the luminous flame ? 10. Put a bit of 
 antimony upon a clean fragment of porcelain and laying it 
 
LABORAl'ORY EXERCISES. 5 I 
 
 upon charcoal or holding it with the test tube holders direct 
 upon the antimony the extreme tip of the blow-pipe flame. 
 More or less of the yellow-white Sb^Og is produced. 1 1. Now 
 bring the porcelain a little nearer and use the inner flame. 
 The oxygen is removed in patches, leaving a coating of anti- 
 mony, dark and shining. These experiments illustrate oxi- 
 dizing and reducing flames. 12. Blow from a dry glass tube, 
 at a distance of a few inches, about ic.c. of lycopodium upon 
 the lamp flame. Dust explosion. 
 
 Paragraphs 43-48 have reference to the formation and 
 some of the properties of the more familiar metallic com- 
 pounds. . In all cases the bearing of the references upon the 
 subject of the paragraph should be written in the notes, 
 though the experiments referred to need not be repeated. 
 
 Oasids (Ctnnpounds with O). 
 
 43. Formation/ (a) By direct combination, i. Light one 
 end of about 3 in. of magnesium ribbon. Collect the product 
 {white MgO) in a clean evaporating dish. 2. Add a little 
 water ; notice the effect on litmus. 3. Upon a fragment of 
 porcelain, held as in 41, 10, heat a little piece of cadmium 
 steadily in the outer blow-pipe flame. Brown CdO is produced. 
 Recall 8, S, IT, 6, 21, 39, 5, 33, 5, 41, 10. (d) By heat- 
 ing certain salts : 4. Heat in a test tube i gm. of copper car- 
 bonate (CuCOg) ; dark brown CuO is left ; what gas escapes ? 
 Prove it. 5. Heat in a glass tube closed at one end J gm. of 
 mercuric nitrate (Hg(N03)2) ; red HgO is left ; what is the 
 result of continuing to heat the HgO ? See 6, i. 6. Heat 
 in a test tube i gm. of ferrous oxalate. Notice how the hot 
 FeO when poured out becomes red hot and changes to a rust 
 color. It oxidizes to FegOj. What gases escape } Suppose 
 
52 EXERCISES IN CHEMISTRY. 
 
 you collect a test tube full over dilute NaOH, which gas 
 would go into the test tube? Recall also 35 and 36, 3. 
 (c) From the hydroxid : 7. To solution of CuSO^ add a few 
 drops of NaOH solution ; pale blue copper hydroxid, chang- 
 ing to dark CuO on warming. 8. To solution of FejClg add 
 NH4OH to alkaline reaction ; filter and heat some of the 
 precipitate on a fragment of porcelain ; red brown FcjOg is 
 left. Recall 7, 8, 33, 2, 37, 2. 
 
 Equations: (i) Mg + 0= . (2) MgO + H20= . (3) 
 
 Cd + O = . (4) CuCOs = + ■ (5) Complete reaction 
 
 uncertain; the brown fumes are oxids of Nitrogen. (6) PeCs04=Pe0 4" 
 
 + (7) CUSO4 + ? NaOH = CuCHa -h . CuOjH, 
 
 = CuO + (8) Pe^Cle + ? NH4OH = Pe,0„H„ + ? NH.Ca. 
 
 Pe^OoHo = + 
 
 JSydroadds {Compounds with OH). 
 
 43. Recall 9, 7, 43, 2, 33, 3 and 5, 43, 7 and 8, 
 also 11. 
 
 Sulfids (^Compounds with S). 
 
 44. Recall 5, 2, 6, 2, 18, i, 4, 6, 34, i. Put a little 
 SbgSg upon charcoal and heat it with the oxidizing flame ; 
 notice the odor of SOj and the formation of some whitish 
 SbjOg. Recall 11, 4, n, 8, 38, 4. 34, 3. 
 
 Chlorids ^Compounds with CI). 
 
 45. Recall 7, 2, 3, 9, 8, i, 13, i, 36, i, 39, i, 2. 
 
 1. To a solution of Pb(N03)2 or lead acetate add dilute 
 HCl, drop by drop, till no more precipitate (PbClj)" is formed. 
 
 2. The same as i, using a solution of any soluble chlorid 
 (NaCl or NH^Cl) instead of HCl. This method applies to 
 
LABORATORY EXERCISES. 53 
 
 silver and mercurous salts also, and distinguishes them with 
 lead from other metals. [3. Arrange for the preparation of 
 H, using dilute HgSO^ and Zn in either a s. n. test tube or 
 an Erlenmeyer flask. But look over 7, and remember par- 
 ticularly about knowing the H is pure before applying a 
 light. Put in a piece of glass tubing about \ gm. of FeClg 
 (ferrous chlorid), connect with the H generator, heat the 
 tube under the FeClg; hold moistened litmus paper at the 
 mouth of the tube ; this proves an acid. Dip a clean glass 
 rod in AgNOg solution and hold at the mouth of the tube ; 
 HCl, if present, will form white curdy AgCl. Compare with 
 7, 8.] Recall 14, i, 4. 16, 4- 
 
 Equations: (l) PbCNOOs + 2HC1 = + . (2) PbCNOs)^ 
 
 + 2NaCl= + (3) PeCl, +H, = Fe+ ■ AgNO. . 
 
 + HCa = + 
 
 Nitrates {Compounds with JVO3). 
 
 46. Recall 36, i, 38, i. i. Dissolve 2 gm. PbCOg in 
 as little hot dilute HNO3 as will dissolve it; Pb(N03)2 is 
 formed which crystallizes from a strong solution on cooling. 
 2. Sprinkle a little KNO3 upon glowing charcoal ; deflagra- 
 tion ; it is an oxidizing agent. Recall 35, 43, S- 
 
 Equation: PbCO3 + 2HNO3=PbCNO0.i+ + 
 
 Sulfates (Compounds with SO 4). 
 
 47, Look up carefully the references under 30; see 
 how many sulfates you find and how they were produced. 
 To a solution of Pb(N03)2 add dilute HjSO^ as long as it 
 gives a precipitate — PbSO^. By a similar method could be 
 made BaSO^, SrSO^, and less easily CaSO^. 
 
 Equation : Pb(NO0 i + H.,SOi = PbSOi -f- 
 
54 EXERCISES IN CHEMISTRY. 
 
 Carbonates iCompounds with COz), 
 
 48. Recall 39, 43, 4. To a solution of BaCl2 add solu- 
 tion of (NH^)2C03. White precipitate of BaCOg. This 
 method applies to many of the metals, but in some cases 
 the carbonate of the metal soon changes to the hydroxid; 
 e. g. iron. See whether the BaCOg will dissolve in HCl — 
 distinction from sulfates (30, 4, S). What gas always 
 escapes when a carbonate dissolves in an acid ? 
 
 Equation: BaCl^ + (NH4)2C03 = BaCOj + 
 
 The Formation of Insoluble Compounds — Precipitation. 
 
 49. This occurs upon mixing soluble substances, when 
 they, by the exchange of their corresponding parts, can form 
 a compound which is insoluble. The corresponding parts 
 are metals, or acid radicals. We have had many illustrations 
 in the foregoing paragraphs ; e. g. 6, 3, 8, 2, IT, 8, 18, 
 4, 6, 30, 4. 5, 34, I, 33, i, 2, 3, 35, 2, 36, i, 43, 7. 8, 
 45, I, 47, 48. In the following experiments (i) note 
 the soluble compounds given; (2) note the possible products, 
 by double decomposition — exchange of corresponding parts; 
 (3) consult Table of Solubility to ascertain whether either of 
 the products is insoluble ; (4) make prediction ; (5) finally 
 mix the substances and see, whether their action corresponds 
 to your prediction. Thus (i) lead nitrate and potassium iodid 
 (2) might give potassium nitrate and lead iodid ; (3) potassium 
 nitrate is soluble, lead iodid is (nearly) insoluble ; (4) there 
 will be then a precipitate of lead iodid, while potassium 
 nitrate remains in solution. (5) Try them. Will there be a 
 precipitate upon mixing solutions of i. Pb(N03)2 and KI .' 
 2. NagSO^ and KI .? 3. BaCl^ and Na^SO^ .? 4. CuSO^ and 
 NagSOi 1 S. CuSO^ and NaCl ? 6. HajCOg and Pb(N03)a ? 
 
LABORATORY EXERCISES. 55 
 
 Table of Solubility for 49. 
 
 Carbonate. 
 
 Chlorid. 
 
 lodid. 
 
 Nitrate. 
 
 Sulfate. 
 
 Barium 
 
 I 
 
 S 
 
 S 
 
 S 
 
 I 
 
 Copper (ic) 
 
 I 
 
 s 
 
 S 
 
 S 
 
 s 
 
 Lead 
 
 I 
 
 I* 
 
 I* 
 
 S 
 
 I 
 
 Potassium 
 
 S 
 
 s 
 
 s 
 
 s 
 
 s 
 
 Sodium 
 
 S 
 
 s 
 
 s 
 
 s 
 
 s 
 
 Solubility and insolubility are of great importance in 
 chemical operations because: — (i) Substances are often 
 separated from one another by the application of solvents, 
 e. g. salt from silicic oxid by water, in 34, 3. (2) Sub- 
 stances are often made, in the laboratory and on the manu- 
 facturing scale, by precipitation ; see Part IV, 8. (3) 
 Qualitative analysis depends largely upon a knowledge of 
 the unlike behavior of different substances with precipitating 
 and dissolving reagents ; e. g. how could you distinguish a 
 solution of either a silver, lead or mercurous salt from other 
 metals ? and also how could you distinguish a salt of lead 
 from either of the other two .■" See last part of this section. 
 For another illustration see 18, 6. (4) Quantitative analysis 
 oftem proceeds by converting substances into insoluble com- 
 pounds and weighing ; e. g. all of the chlorin in a solution of 
 HCl may be combined by the addition of AgNOg, into 
 insoluble AgCl, from the weight of which the CI and the 
 HCl may readily be calculated. 
 
 It may be convenient for the learner to remember that, 
 speaking generally, the following compounds are soluble in 
 ■water: compounds of sodium, potassium, ammonium ; ni- 
 trates ; acetates ; chlorids (except of Ag, HJ, Pb) ; sulfates 
 (except of Pb, Ba, Sr, Ca) ; bi or acid carbonates. Generally 
 
 * Soluble in hot water, or large quantities of cold water. 
 
56 EXERCISES IN CHEMISTRY. 
 
 insoluble are metals and oxids of metals (except in a few 
 cases with chemical change) ; hydroxids ; sulfids ; normal 
 carbonates ; phosphates ; (except compounds of Na, K, NH^ 
 and except the hydroxids and sulfids of Mg, Ca, Sr, Ba). 
 Each compound, however, has its own peculiar solubility, 
 which can only be known by experience and consultation of 
 the more complete tables of solubility. 
 
 Action of Metals upon Solutions of Metallic Salts. 
 
 [50. In the following experiments two facts are to be 
 verified in each case — first, that the metal in solution is 
 partly or wholly deposited, evidenced by sight of the 
 deposit; second, that the metal at the outset uncombined 
 has partly entered into the solution, shown by testing the 
 solution with some characteristic reagent before the change, 
 and then testing it, or another portion of it, after the 
 change. 
 
 I. Into each of two test tubes put about 8c.c. of AgNOg 
 solution ; to one tube add a globule of Hg ; HgNOg and ^Ag 
 are produced. After five to ten minutes look for the Ag, 
 then add KI to both tubes. Agl is yellow, Hgl is greenish. 
 If any Hgl is produced, it shows that Hg had gone into the 
 liquid as a soluble compound. 2. Similarly HgCl2+Cu= 
 CuClg+Hg. Test with NH^OH ; a blue liquid shows copper. 
 3. Cu(N03)2+Pb=Pb(N03)2+Cu. Test with dilute H2SO4; 
 Pb will give white PbSO^. 4. Pb(N03)2+Zn=Zn(N03)2-l- 
 Pb. Test with K3FeCgNg (ferricyanid of potassium) ; 
 ferricyanid of Zn is brownish. 
 
 Notice that Zn displaces Pb ; Pb displaces Cu ; Cu dis- 
 places Hg ; Hg displaces Ag.] 
 
LABORATORY EXERCISES. 57 
 
 To the Pupil. 
 
 At the conclusion of this part of your work you will be 
 expected to identify the substances used ; at least those 
 included in the following list. You will have no difficulty 
 in doing this, if you follow this plan : — 
 
 1. Define. 
 
 2. Classify. 
 
 3. Distinguish. 
 
 1. Define, by writing against each substance a very brief 
 description — viz. "a blue liquid," "a yellow, brittle solid." 
 
 2. Classify, by arranging in groups according to obvious 
 properties ; this can be done in many ways — solids, liquids ; 
 colored, not colored ; with odor, without odor ; affecting 
 litmus, not affecting litmus ; metals, not metals, etc. 
 
 3. Distinguish. If you classify well, the groups distin- 
 guish themselves, and you have only to find some physical 
 or chemical difference between a few of the same group. 
 Thus you can easily tell copper from iron or zinc or cadmium 
 or tin by its appearance ; you can tell cadmium from tin, not 
 by its appearance but by its brown oxid when heated with 
 the blow-pipe or by the fact that it dissolves in HNO3, while 
 tin does not. 
 
 List. Properties. 
 
 copper sulfate, crystals 
 copper sulfate, solution 
 ferric oxid 
 sulfur, broken roll 
 sulfur, flowers 
 copper, foil 
 copper, sheet 
 iron wire 
 
58 EXERCISES IN CHEMISTRY. 
 
 List. Properties. 
 
 mercuric oxid 
 zinc dust 
 zinc, granulated 
 hydrochloric acid, cone, 
 hydrochloric acid, dil. 
 sulfuric acid, cone, 
 sulfuric acid, dil. 
 potassic chlorate 
 sugar 
 starch 
 lime 
 
 lead oxid 
 
 carbonate of soda, cryst. 
 calcic chlorid, fused 
 " chlorid of lime " 
 hydrogen peroxid 
 permanganate of potassium, sol. 
 sulfate of lime 
 common salt, fine 
 rock salt 
 
 bromid of potassium 
 iron sulfid 
 iodin 
 
 iodid of potassium 
 paraffin 
 water 
 tin 
 
 cadmium 
 baric chlorid, sol. 
 lead carbonate 
 silver nitrate, sol. 
 manganese dioxid 
 ferric chlorid, sol. 
 alcohol 
 acetic acid 
 sulfate of soda, sol. 
 lead nitrate, sol, 
 sulfite of soda 
 ammonium chlorid 
 
LABORATORY EXERCISES. 59 
 
 List. Propertiss. 
 
 lime water 
 ammonia, sol. 
 ammonium nitrate 
 nitric acid, cone, 
 nitric acid, dil. 
 potassimn nitrate 
 bisuliid of carbon 
 antimony 
 sulfid of antimony 
 silicate of soda, sol. 
 charcoal 
 bone black 
 carbonate of lime 
 caoutchouc 
 lycopodium 
 magnesium 
 carbonate of copper 
 iron oxalate 
 
 How could you distinguish copper from charcoal ? 
 
 How could you distinguish solution of NH^Cl from solu- 
 tion of NaCl ? 
 
 How could you distinguish solution of NaCl from NagSO^ ? 
 
 How could you distinguish fine salt from ferrous sulfid ? 
 
 How could you distinguish fine salt from lead carbonate ? 
 
 How could you distinguish dilute HCl from HgSO^.'' 
 
 How could you distinguish bisulfid of carbon from any- 
 thing else .' 
 
 How could you distinguish O from H .' 
 
 How could you distinguish N from H > 
 
 How could you distinguish HCl gas from NHg gas .' 
 
 How could you distinguish SOj from HjS .' 
 
 How could you distinguish CaCOg from NajCOj ? 
 
PART II. 
 
 CHEMICAL COMPUTATIONS 
 
 Atomic weights. 
 
 Class I. Constants from experimental data 
 
 1. Vapor Density. 
 
 2. Molecular Weight of Volatile Substances. 
 
 3. Molecular Weight of Solids. 
 
 4. Atomic Weight. 
 
 5. {a) Formula of Volatile Substances. 
 {b) Formula of Solids. 
 
 6. Valence. 
 
 7. Equations. 
 
 Class II. Practical calculations from accepted constants. 
 
 8. Molecular or Formula Weight. 
 
 9. Vapor Density. 
 
 10. Weight of gas, volume given. 
 
 11. Volume of gas, weight given. 
 
 12. Correction of gaseous volumes. 
 {a) For Temperature. 
 
 {b) For Pressure. 
 (c) For Tension. 
 
 13. Percentage, from Formula. 
 
 14. Weight of constituent of compound. 
 
 15. Computations from Equation. 
 (a) Weight. 
 
 (i^) Volume. 
 
 16. In the Metric System. 
 
 {a) Weight of liquids from measure. 
 {b) Measure of liquids from weight. 
 
 17. Normal Solutions. 
 Examples applying these principles. 
 
CHEMICAL COMPUTATIONS. 
 
 6i 
 
 A Table of Atomic Weights (0=i6). 
 ary i, 1895, by F. W. Clarke. 
 
 Revised to Janu- 
 
 NAME. SYMBOL. ATOMIC WEIGHT. 
 
 Aluminum Al 27. 
 
 Antimony Sb 1 20. 
 
 Arsenic As 75. 
 
 Barium Ba 137-43 
 
 Bismuth Bi 208. 
 
 Boron B II. 
 
 Bromin Br 79-95 
 
 Cadmium Cd 112. 
 
 Caesium Cs 132.9 
 
 Calcium Ca 40. 
 
 Carbon C 12. 
 
 Cerium Ce 140.2 
 
 CHorin CI 35.45 
 
 Chromium Cr 52.1 
 
 Cobalt .'Co 59.5 
 
 Columbium Cb 94. 
 
 Copper Cu 63.6 
 
 Erbium Er 166.3 
 
 Fluorin F 19. 
 
 Gadolinium Gd 156.1 
 
 Gallium Ga 69. 
 
 Germanium Ge 72.3 
 
 Glucinum GI 9. 
 
 Gold Au 197-3 
 
 Hydrogen H 1.008 
 
 Indiimi In i'3.7 
 
 lodin I 126.85 
 
 Iridium Ir i93-i 
 
 Iron Fe 56. 
 
 Lanthanum La 13S.2 
 
 Lead Pb 206.95 
 
 Lithium Li 7.02 
 
 Magnesium Mg 24.3 
 
 Manganese Mn 55. 
 
 Mercury Hg 200. 
 
 Molybdenum . . • .Mo 96. 
 
 ATOMIC WEIGHT. 
 
 Neodymium Nd . 
 
 Nickel Ni . . 
 
 Nitrogen N . . 
 
 Osmium . . • ■ Os • 
 
 Oxygen O . . 
 
 Palladium Pd . 
 
 Phosphorus P . . . 
 
 Platinum Pt . 
 
 Potassium K . . 
 
 Praseodymium . . .Pr. . 
 
 Rhodium Rh . 
 
 Rubidium Rb . 
 
 Ruthenium Ru . 
 
 Samarium ....... Sm . 
 
 Scandium Sc . 
 
 Selenium Se . 
 
 Silicon Si ■ . 
 
 Silver Ag . 
 
 Sodium Na . 
 
 Strontium Sr . . 
 
 Sulfur S... 
 
 Tantalum Ta . 
 
 Tellurium Te . 
 
 Terbium Tb . 
 
 Thallium Tl .. 
 
 Thorium Th - 
 
 Thulium Tu . 
 
 Tin Sn . 
 
 Titanium Ti . . 
 
 Tungsten "W . 
 
 Uranium U . . 
 
 Vanadium V . . 
 
 Ytterbium Yb. 
 
 Yttrium Yt . 
 
 Zinc Zn . 
 
 Zirconium Zr.. 
 
 140.5 
 58.7 
 14.03 
 
 190.8 
 16. 
 
 106.5 
 31- 
 
 195- 
 39" 
 
 I43-S 
 
 103. 
 
 85-5 
 101.6 
 150. 
 
 44. 
 
 79- 
 28.4 
 
 107.92 
 23.05 
 87.66 
 32.06 
 
 182.6 
 
 125. 
 
 160. 
 
 204.18 
 
 232.6 
 
 170.7 
 
 119. 
 48. 
 
 184.9 
 
 239.6 
 51-4 
 
 173- 
 89.1 
 
 65-3 
 90.6 
 
62 EXERCISES IN CHEMISTRY. 
 
 CHEMICAL COMPUTATIONS. 
 
 Class I. — Based upon experimental data, and illustrating 
 methods of establishing Atomic Weights, Formulas, and 
 Equations. 
 
 1. The vapor density of a gas is its weight in compari- 
 son with the weight of its own volume of hydrogen. 
 
 Note. — Chemists have generally accepted the view stated in the "Theory of 
 Avogadro"; viz. that equal volumes of gases contain the same number of 
 molecules. According to this view the number obtained for Density as above 
 would also be the weight of I molecule (of the gas under discussion) in com- 
 parison with 1 molecule of hydrogen. 
 
 Method : Weighing a definite volume of the gas. 
 
 Illustration: The weight of i L of ammonia gas is .7616 gm.; the weight 
 of : L of hydrogen is .0896 gm. .7616 -^ .0896 = 8.5. The vapor density of 
 ammonia is 8.5, i. e. it is 8.5 times as heavy as the same volume of hydrogen 
 under standard conditions . 
 
 3. The molecular weight of a gas is its weight compared 
 with the weight of half its own volume of hydrogen — i. e. 
 twice its density. 
 
 Note. — According to Avogadro's Theory, molecular weight, as obtained above, 
 would be the weight of i molecule (of the gas under discussion) in comparison 
 with the half-molecule of hydrogen. 
 
 Datum : Vapor density. 
 
 Method : Multiplication by 2. When the density is for any reason stated in 
 terms of air, the multiplication is by 14.44 X 2. 
 
 Illustration: The vapor density of O, N, CO2 is 16, 14, 22, respectively; 
 the molecular weight is 32, 28, 44. 
 
 3. The molecular weight of a non-volatile solid is in 
 some cases determined by certain physical methods, for 
 which reference must be made to a larger treatise. 
 
 E. G. — See " Solutions" by W. Ostwald, upon Vapor 
 Pressures and Freezing Points, 
 
CHEMICAL COMPUTATIONS. 63 
 
 4. The atomic weight of an element is the relative 
 weight (*H being i or O being i6) in which the element 
 enteKs into chemical changes and forms part of chemical 
 compounds. It is determined by as many of the following 
 methods as can be applied : — 
 
 A. Data : Percentage composition and vapor density of 
 a large number of gaseous compounds in which the element 
 occurs. 
 
 Method : The vapor densities give the molecular weights of the compounds; 
 the percentages give the quantity of the element, whose atomic weight is sought, 
 in those molecular weights. The greatest common divisor of the numbers 
 expressing the quantity of the element is the atomic weight. 
 
 Illustration : Oxygen occurs in water, carbon dioxid, sulfur trioxid, etc. 
 Their vapdr densities are 9, 22, 40, and their molecular weights, 18, 44, 80, 
 respectively. Oxygen is 88.89% of the water, or \%; 72.72% of carbon dioxid, 
 or fl; and 60% of sulfur trioxid, or |^. 16 is considered the atomic weight 
 of oxygen. 
 
 B. Data : The weight of the element which displaces i 
 part of hydrogen from a compound. 
 
 C. Data : The weight of the element which combines 
 with 35.45 parts of .chlorin. 
 
 Method : The number obtained by either B pr C is the same, and is either 
 the atomic weight, or a sub-multiple of the atomic weight. This method is 
 accurate for the ratio. 
 
 Illustration : 107.9 parts of silver displace i of hydrogen. 103.47 °f l^^d 
 combine with 35.45 of chlorin. The atomic weights of silver and lead are 
 either 107.9 and 103.47, or two or three, etc., times these numbers. See D. 
 
 * Note. Since the ratio of O to H is not at present considered exactly 16 
 
 to I, but 15.88 to I, two slightly differing systems are in use — one in which the 
 computations are to O as 16 (H = 1.008), and the other in which H is taken as 
 I (O = 15.88). In the table on page 61 O = 16 is the standard. 
 
64 EXERCISES IN CHEMISTRY. 
 
 D. Data : The specific Reat of the element, and the 
 number 6.4*, which is called the atomic heat, and which is 
 approximately the same for all solid elements. 
 
 Method : 6.4 divided by the specific heat of the element gives its atomic 
 weight. This result is approximate only, but settles the question left undecided 
 by B and C. 
 
 Illustration: The specific heat of silver is .0559. 6.4^ .0559 = 'H! ^nd 
 the atomic weight of silver is 107.9, and not a multiple of 107.9. The specific 
 heat of lead is .0315. 6.4 -:- .0315 = 203; and the atomic weight of lead is not 
 103.47, '^i' twice as much, or 206.95. 
 
 5. A. A chemical formula of a gas or volatile substance 
 expresses the proportional composition of the substance and 
 a total weight equaling twice its vapor density. 
 
 Data : The vapor density of the compound, the proportions of the constituent 
 elements in the compound, and the atomic weights of the elements. 
 
 Method : Obtain molecular weight : compute the parts of it indicated by the 
 proportion of each element in the compound: these numbers divided by the 
 atomic weights of the corresponding elements give the number of atoms of each 
 element. 
 
 Illustration : Water has a vapor density of 9 and molecular weight of 18, 
 and consists of 11.14% hydrogen and 88.86% oxygen. The atomic weights of H 
 and O are i and 16 respectively. 11.14% of 18 = 2 -:- I = 2 or Hj. 88.86% of 
 18 = 16 -f- 16 = I or O. Formula, HjO. 
 
 B. A chemical formula of a non-volatile substance usually 
 expresses the proportional composition of the substance in 
 the simplest manner consistent with the weight meaning 
 implied in the symbols. 
 
 Data : Proportional composition of the substance, and the atomic weights of 
 the elements. 
 
 Method : Divide the percentage numbers by the atomic weights respectively : 
 results represent relative number of atoms : these results are usually fractional, 
 and are to be reduced to integers. 
 
 * Variously estimated; — from 6.1 to 6.5. 
 
CHEMICAL COMPUTATIONS. 65 
 
 Illustration : 
 
 Analysis gives, C 35.82 , -5- 12 = 2.985, -i- 2.985 = 1 , X 4 = 4. 
 H 4-479. -5- I = 4-479. "^ z-98S = 4, X 4 = 6. 
 O 59.7 , -T- 16 = 3.731, -5- 2.985 = 4, X 4 = 5- 
 
 99-99+ 
 Formula is C4H6O11. 
 (This method is now being modified by the new methods for determining mo- 
 lecular weight of solids suggested in 3. ) 
 
 6. The Valence (Equivalence, Quantivalence) of an 
 element is practically the number of times i of H, 35.4 of 
 CI, or 8 of O, with whi(?h its own atomic weight unites. 
 
 The idea of valence depends upon the correct determina- 
 tion of formulas and atomic weights. Thus from the 
 formula HCl — i to 35.4 — we say the valence of CI is i. 
 Its atomic weight combines with the atomic weight of the 
 standard — H. From the formula NaCl — 23 to 35.4 — we 
 say the valence of Na is i ; from the formula BaClj — 137 
 to 70.8 — we say the valence of Ba is 2. SnO and SnOj 
 seem to show the valence of Sn as sometimes 2, sometimes 
 4. Valence is indicated — H', CI', Ba", Sn"- " 
 
 Valence is useful in assisting .the memory of formulas, or 
 in conjecturally writing unknown formulas ; but it must be 
 remembered that, experimentally, valence is an inference 
 from formulas, not formulas from valence. 
 
 7. A chemical equation indicates, in their relative pro- 
 portions, all the materials which enter into and proceed 
 from a chemical change. 
 
 Data: The only complete data for an equation are the weights of all the 
 factors and products. When, however, all the factors and products are known, 
 and their formulas, a correct equation can usually be written. 
 
 A. Method : Inspection. Write the smallest number of whole molecules 
 which balance. 
 
66 EXERCISES IN CHEMISTRY. 
 
 Illustration: SbCl^ + H„S = Sb^Ss + HCl. 
 
 2SbCl3 + 3H,S = Sb,Sa + 6HC1. 
 B. Method : Assume coefficients, and form equations between the coefficients 
 of the same element. Solve. 
 
 ILLUSTKATION : NijOs + HCl = NiCU + H^O + CI. 
 
 sNiiiOs + tHCl = xNiCU + yHaO + zCl. 
 2s = X, 3s ^y, t = 2y, t = "2x + z. Assume s =: 1, then x = 2, y = 3, 
 t =6, z = 2, and the equation is Ni^O^ + 6HC1 = 2NiCl2 + 3H2O + 2C1, or CI2. 
 
 Class II. — Practical calculations in which Atomic 
 Weights, Formulas, and Equations are accepted as estab- 
 lished facts. ^ 
 
 8. To compute molecular weight or formula weight. 
 
 Data: The formula of the compound; atomic weights, by implication. 
 Method : The sum of the atomic weights in the formula is the molecular or 
 formula weight. 
 
 Illustration : The molecular weight of HjSO, = 2 + 32 + 64 = 98. 
 
 9. To compute vapor density (of a gas or volatile 
 substance). 
 
 Data: Formula. 
 
 Method : Take half the molecular weight. 
 
 Illustration : The density of afcohol vapor (C2H5OH) is 
 
 24 + 5 + 16 + 1 _, 3 
 2 
 This gives hydrogen density; if air density is desired, the result is divided by 
 14.44. 
 
 10. To compute the weight of a given volume of gas. 
 
 Data : Formula or density of gas, and number of litres; also the weight of 
 I L. of hydrogen (.0896 grammes). 
 
 Method : Multiply the density of the gas by .0896 gr., and the product by 
 the volume of the gas in litres. W = V. D. X .0896 X L. 
 
 Illustration : The weight of 25 L. of oxygen, V. D. = 16, is 16 X .0896 .X 
 35 = 35.84 gr. 
 
CHEMICAL COMPUTATIONS. 6"] 
 
 11. To compute the volume of a given weight of gas. 
 
 Data : Formula or density of gas and total weight; weight of i L. of H. 
 Method : Divide the given weight by the density times .0896 gr. 
 W 
 
 V. or L. = 
 
 V. D. X .0896 
 
 Illustration : How many litres are 17 grammes of COj? 
 "7 
 
 22 X .0896 
 
 = 8.62 L. 
 
 13. To correct measured volume of gases for tempera- 
 ture, pressure, and tension of water vapor. 
 
 The volume of all gases is referred to 0°, 760m. m. and a 
 condition of dryness. It is not practically convenient to 
 measure under exactly the above conditions ; so gases are 
 measured at convenient temperature and pressure and cor- 
 rections are made. 
 
 A. For Temperature: Multiply the measured volume by 273 and divide 
 by 273 + (or — ) the number of degrees above (or below) 0°C. 
 
 V X 273 
 
 ^ — 273 =tt ■ 
 
 B. For Pressure : Multiply the measured volume by the number of milli- 
 metres of barometric pressure, and divide by 760. 
 
 ^ ~ 760 
 Either correction may be made upon the result of the other, or both may be 
 combined in one operation, 
 
 C. The tension of water vapor is usually aj^iplied from a table and subtracted 
 from P the observed pressure (at I5°C the correction is I2.7m.m.). 
 
 Illustration: A gas measures over water, 450C.C. at 75om.m. and is°C. 
 What is its normal volume ? 
 
 V =450x1x^^=413.8. 
 
68 EXERCISES IN CHEMISTRY. 
 
 13. To compute percentage composition of any com- 
 pound from its formula. 
 
 Data : Formula, implying atomic, weights. 
 
 Method : Take such part of loo as the weight of any constituent, expressed 
 by the formula, is of the whole molecular weight. 
 
 Illustration : The percentage composition expressed by NajCOs is com- 
 puted as follows : — 
 
 Na^ = 46 Nas = -[^ of 100 = 43.39%. 
 C =12 C =-[1^ of 100=11.32%. 
 Os =48 O3 =1^ of 100 = 45.28%. 
 
 106 
 
 14. To compute the weight of any constituent in a 
 given weight of a compound. 
 
 Data : The formula of the compound and the given weight. 
 Method : Take such part of the given weight as the weight of any constit- 
 uent, expressed by the formula, is of the whole molecular weight. 
 
 Illustration : How many grammes of oxygen in 200 grammes of KCIO3 ? 
 
 K = 39. 
 
 Cl = 3S-5. 
 
 O3 = 48. O3 = tM .s °f 200 or 78.36 grammes. 
 
 122.5 
 
 15. To compute from equations. 
 A. Weight. 
 
 Data : The equation and the weight of one of the compounds. 
 
 Method : The weight of any compound is to the weight of any other com- 
 pound as the relative weight in the equation of the first is to the relative weight 
 of the other. W : w = M : m. 
 
 Illustration : NajCOs + HaSOi = Na^SO, + CO2 + HjO. 
 
 What weight of H2SO4 is needed with 15 gr. of NajCOa? io6:98=i5:x 
 = 13.8- 
 
 How much NaQS04 can be made from 15 gr. of Na^COa? 106:142 = 
 15 :x=: 20. 
 
 106, 98, 142 are the relative weights expressed by NajCOs, H2SO4, Na2S04. 
 
CHEMICAL COMPUTATIONS. 69 
 
 B. Volume. 
 
 When the equation involves gases each formula represents 
 the same volume. (This is an inference from 1, 2 and 5). 
 Thus CH^-f- 2O2 = CO2 + 2H2O, and volume computations 
 
 a vols. 4 vols. 2 vols. 
 
 can be made ; but O^, H^, Ng, C\ are the formulas (not O, 
 H, N, Ciy of simple gases. 
 
 Ilxustration : How many litres of O are required to burn I litre of CO ? 
 Equation : 2CO + 0, = 2CO2. 
 
 4 vols. 2 vols. 4 vols. 
 
 4:2::! : X, X = ^. Ans^ 
 
 16. In the metric system of weights and measures the 
 relations between weight, measure, and specific gravity are 
 very convenient, i cubic centimeter of water (by measure) 
 = I gram (by weight) ; and the specific gravity of water is i. 
 
 To calculate weight from measure, of liquids. 
 
 Data : Measure and specific gravity. 
 
 Method : Multiply the number of cubic centimeters by specific gravity : the 
 result is grams. 
 
 Illustration : 400C.C. of alcohol, specific gravity .80, weighs 400 X .80 =: 
 320 grams. 
 
 B. To calculate measure from weight, of liquids. 
 Data : The weight and specific gravity. 
 
 Method : Divide the number of grams by the specific gravity. 
 Illustration: 370 gm. of H2SO4, specific gravity 1.85 measures 370 -j- 1.85 
 = 200C.C. 
 
 IT. Normal solutions. 
 
 A normal solution of a chemical reagent means primarily 
 the formula weight of the substance, in grams, with 
 enough water to measure i litre- (i,oooc.c.). Thus normal 
 sodium hydroxid (NaOH) would contain 40 gm. in i litre of 
 
70 EXERCISES IN CHEMISTRY. 
 
 solution, or 40 mgm. in i cubic centimeter. In the case, 
 however, of diacid bases, like Ca(0H)2, and dibasic acids, like 
 H2SO4, one-half the formula weights, 37 and 49, would be 
 used, thus giving ic.c. the same neutralizing power as ic.c. 
 of NaOH or HCl ; and other reagents are made up on the 
 same principle. 
 
 In practice much more dilute solutions are used — -^ and 
 2^5- normal — which are often written -^ and ^. 
 
 The strength of exact solutions is often indicated frac- 
 tionally, the numerator denoting the weight of reagent in 
 grams, and the denominator the total measure of the 
 solution in centimeters ; thus -jf^ might represejit normal 
 KOH. 
 
 The principles embodied in the problems given above find 
 many applications. A few are given below. The teacher 
 can multiply such problems indefinitely and give them out 
 upon numbered cards. 
 
 The pupil is advised to indicate all the arithmetical work 
 in each operation before performing it ; also, he should, as 
 soon as able, perform multiplications and divisions by means 
 of logarithms. After the principle is understood, the work 
 is much simplified by the use of Tables, e. g. Adriance's 
 Laboratory Calculations. 
 
 1. Calculate the molecular weighis of HCl, HjS, CO, COg, 
 NH3, N2O, SO2. 
 
 2. Calculate the vapor densities of each of the above 
 gases. 
 
 3. Which of them are heavier than air.'' 
 
 4. Calculate the percentage composition of KBr, NaBr, 
 NH^Br, NiBr2, MgBrg, LiBr, and arrange them in the order 
 of most or least Br. 
 
CHEMICAL COMPUTATIONS. /I 
 
 5. Calculate the true volume of 300C.C. of oxygen, 
 measured at 20°C and 74om.m. pressure. 
 
 6. What is the normal volume of 300C.C. of oxygen 
 measured at — 2o°C and /Som.m. .' 
 
 7. Compute the weight of i litre of each of the gases 
 mentioned in i. 
 
 8. If 9 parts of Al combine with 35.45 of CI and the 
 specific heat of Al is .2253, what is the atomic weight of Al ? 
 
 9. What is the per cent, of iron in Fe^ (PjO.jyg, QH^O .' 
 
 10. What weight and volume of H is obtainable by 4 gm. 
 of Zn } 
 
 11. Compute the formula of a gas found by experiment to 
 consist of N, 63.63%, O, 36.36%, and to have vapor 
 density 22. 
 
 12. How many centimeters of HCl solution, specific 
 gravity 1.21, containing 42.4% HCl by weight, are needed 
 for 4 litres of chlorin > 4HC1 + MnOj = Clj, etc. 
 
 13. Assign formula to a solid found by experiment to 
 consist of Ca, 29.41% ; S, 23.52% ; O, 47.05%. 
 
 14. Na= 39.31%; CI =60.68%; assign forniula. 
 
 15. See n, d5. (i) How much HjSO^ with 10 gm. 
 NajCOg.? (2) How much HjSO^ for 10 gm. NagSO^? 
 (3) How much H2SO4 for 10 gm. CO2.'' (4) How much 
 NagCOg with 10 gm. H^SO^.? (5) For 10 gm. Na^SO^? 
 (6) For 10 gm. COj ? (7) How much COg from 10 gm. 
 NagCOg.' (8) How much COg when 10 gm. H2SO4 are 
 used ? (9) When 10 gm. NagSO^ are produced ? (10) How 
 much NagSO^ from 10 gm. HgSO^.? (11) From 10 gm. 
 NagCOg.'' (12) When 10 gm. COg are produced .' 
 
 16. In 2,000 lbs. of NaNOg, 85% pure, how much Na? 
 
72 EXERCISES IN CHEMISTRY. 
 
 17. How much H2SO4 is needed to make 25 litres of COj 
 from CaCOg ? 
 
 18. Suppose a manufacturer of HjSO^ to start with 100 
 lbs. of sulfur and to obtain the full theoretical yield in each 
 case : 
 
 (i) What weight of O is needed to burn the S to SO2 .' 
 
 (2) What additional weight of O to convert the SOj 
 
 to SO3 ? 
 
 (3) What weight of Hfi to convert the SO3 to HgSO^ ? 
 
 (4) And how much H2SO4 is obtained ? 
 
 19. What weight of KCIO3 will yield the oxygen to fill 
 two 300c. c. jars ? 
 
 20. A student is experimentally determining the amount 
 of Ba in i gm. of barium chlorid (BaClg, SHjO). He does 
 this by converting all the Ba into BaS04, and weighing 
 that. He obtains .96 gm. BaSO^. How great is his error ; 
 i. e., the difference between the theoretical % of Ba in 
 the barium chlorid and the % computed from his result ? 
 
 21. iic.c. of a r^Q normal solution of nitrate of silver 
 (AgNOg) is just sufficient for all the salt (NaCl) in 250C.C. 
 of a water which is being analyzed. How much NaCl is 
 there in a litre of the water, and how much CI .■' AgNOg + 
 NaCl = AgCl + NaNOg. 
 
 22. A substance believed to be cupric hydrid gives from 
 analysis, copper 96.72%, hydrogen 3.28% ; compare these 
 results with the theoretical results from the formula CuHj. 
 
 23. A solution of HgSO^ intended to be ^^ is found to 
 contain .00524 gm. per cubic centimeter of acid ; how much 
 water must be added to i litre of this solution in order to 
 make the resulting solution exactly -j^.' 
 
PART III. 
 
 EXPERIMENTS IN QUANTITATIVE RELATIONS. 
 
 The Supplies. 
 
 1. The hydrogen set free by a definite weight of metal. 
 
 {a) Zinc and hydrochloric acid. 
 
 (b) Magnesium and sulfuric acid. 
 
 (c) Tin and hydrochloric acid. 
 
 2. The percentage of oxygen in chlorate of potassium ; 
 the volume of a given weight of oxygen. 
 
 3. The percentage of water of crystallization in 
 
 {a) Baric chlorid. ^ 
 
 {b) Copper sulfate. 
 (c) Sodium carbonate. 
 
 4. The weight of salt in solution at a given temperature. 
 
 5. The solubility of a salt at different temperatures. 
 
 6. The proportion by volume of oxygen in air. 
 
 7. The behavior of indicators. 
 
 8. The neutralizing power of an aikali ; a certain weight, 
 and twice that weight. 
 
 (rt) With hydrochloric acid. 
 (b) With sulfuric acid. 
 
 9. The relative neutralizing power of the hydroxids of 
 potassium and sodium. 
 
 10. {a) The quantity of sulfate, chlorid, and nitrate of 
 sodium from a given weight of the carbonate. 
 
 {b) The calculated quantity of acids used. 
 {c) The basicity of the acids. 
 
74 EXERCISES IN CHEMISTRY. 
 
 11. {a) The weight of copper precipitated by cadmium. 
 {b) The weight of lead precipitated by cadmium. 
 
 {c and d) The method applied to silver and mercury. 
 
 12. One-half the volume of hydrochloric acid gas is hydrogen. 
 
 13. The relative volumes of carbon monoxid and dioxid 
 from the decomposition of ferrous oxalate. 
 
 14. Successive transformations of a lead compound with 
 the weights of the products. 
 
 {a) The buff oxid from white lead. 
 
 if)) The yellow oxid from the buff oxid. 
 
 (c) The nitrate from the yellow oxid. 
 
 (d) The sulfate from the nitrate. 
 Summary. 
 
 The experiments in Part III have been selected or devised 
 with reference to reasonable accuracy without elaborate 
 appliances. The pupil's own skill and good sense are 
 important factors. He will commonly improve his work by 
 trying again. Primary emphasis is laid upon the quan- 
 titative relations ; but the pupil must comprehend every 
 chemical change involved, and the reason for every process. 
 
 Each student will require the following articles ; but it is 
 not necessary to get a complete outfit for each of several 
 pursuing the same course. 
 
 I Stoddard's balance, or other balance weighing to 10 m.gr. 
 
 I 50OC.C. flat-bottomed flask. 
 
 I 2-hole rubber stopper for flask. 
 
 I bulbed ignition tube, 4'' x \". 
 
 1 l-hole rubber, or best velvet cork stopper for ignition tube. 
 
 2 royal berlin porcelain crucibles, 20c.c., or 
 
 2 royal berlin porcelain evaporating dishes, 30 to 50C.C. 
 
 I dessicator (a covered dish with a platform conveniently of perforated tin; the 
 
 dish containing about \ in. of concentrated H2SO4). 
 I 2-inch watch glass. 
 
EXPERIMENTS IN QUANTITATIVE RELATIONS. 
 
 75 
 
 I water bath (usually of copper; can be made by fitting a tin cover with suitable 
 
 holes to an ordinary tin pan) . 
 I 50C.C. burette, graduated to ^f,, 
 I stand for burette. 
 I pipette to deliver loc.c. 
 
 I 50OC.C. side-necked Erlenmeyer filtering flask of stout glass. 
 I solid rubber stopper ifor flask. 
 
 I spherical separating funnel \ or J" tube; soft glass tubing, \ or f ". 
 I or 2 stout pinch-cocks; rubber tubing, glass tubing, test tubes, etc., as in Part I. 
 
 The only chemicals required beyond the supply for Part 
 I are indicators : litmus, cochineal, lacmoid, phenol-phthal- 
 ein, etc., magnesium powder, pure, dry, white lead. 
 
 1. Tke quantity of hydrogen set free from acids by definite 
 weights of metals. 
 
 * Apparatus. (See Fig. 11.) It consists of a sooc.c. 
 flask with a 2-hole rubber stopper, a side-necked test tube 
 6" X i" with the neck 
 bent down at a right 
 angle and put through 
 one hole in the stopper 
 of the flask, and rubber 
 stopper to iit ; a gas jar 
 holding about 40OC.C. ; 
 a glass tube bent at right 
 angles passes through 
 the other hole of the 
 stopper almost to the 
 bottom of the flask and 
 is connected to rubber 
 tubing reaching to the 
 bottom of the gas jar. At the beginning of the experiment 
 the flask is filled with water to one inch of the stopper; 
 * Modified from Ramsey's " Chemical Theory." 
 
 Fig. II. 
 
"](> EXERCISES IN CHEMISTRY. 
 
 acid of the specified strength is put in the s. n. test tube, 
 making it about half full ; all the stoppers and joints must 
 be tight. The pupil should do two of either A, B, or C. 
 
 A. Zinc and HCl. Use dilute HCl (i to 2 or 3), weigh 
 out about I gm. of granulated zinc (not less than .9 gm. nor 
 more than i.i). Remove the stopper from the s. n. tube, 
 drop in the zinc and instantly replace stopper. The hydro- 
 gen evolved will now force its volume of water from the 
 flask into the gas jar. When all action is over, and the flask 
 and all the apparatus have cooled to the temperature of the 
 room, bring the liquids to the same level in flask and gas 
 jar, detach rubber tube and let the water in it flow into the 
 jar. Observe temperature of room, and barometric pressure, 
 if possible. Now measure carefully the contents of the gas 
 jar, making the best allowance you can for the glass tube in 
 the flask. 
 
 Data and Computations. 
 
 Number of c.c. of H produced by gm. of Zn = 
 
 Corrected for temperature ( °), pressure ( , m.m.), and tension = 
 
 Weight of H (computed; i litre =: .0896 gm.) = 
 
 Weight of Zn required for I. of H, computed"! 
 
 Equivalent of Zn / 
 
 Accepted equivalent of Zn (= atomic veeight -J- 2) = 
 
 B. Magnesium ribbon (.3 gm.) and H2SO4 (l to 10). 
 
 Data and Computations. 
 
 Number of c.c. of H produced by gm. of Mg ^ 
 
 Corrected for z= 
 
 Weight of H (computed) =; 
 
 Weight of Mg for i of H 1 _ 
 Equivalent of Mg I , 
 
 Accepted equivalent of Mg =: 
 
EXPERIMENTS IN QUANTITATIVE RELATIONS. 
 
 n 
 
 C. Tin (i.s to 2 gm.) and HCl (concentrated). 
 
 Data and Computations. 
 
 Number of c.c. of H produced by 
 
 Corrected for 
 
 Weight of H (computed) 
 
 Weight of Sn for i of H 1 
 Equivalent of Sn J 
 
 Accepted equivalent of Sn 
 
 gm. of Sn : 
 
 3. The percentage of oxygen in KCIO^ 
 of a given weight of O. 
 
 also the volume 
 
 Apparatus : The same as that used in i, except that a 
 clean, dry, bulbed ignition tube is used instead of the s. n. 
 test tube. The ignition tube has a cork stopper bored for 
 glass tubing, which just 
 passes through the cork 
 and the rubber stopper of 
 the flask ; the ignition 
 tube is arranged to be 
 nearly horizontal. (Fig. 12.) 
 Weigh the ignition tube, 
 put into the bulb between 
 I and 1.5 gm. of KClOg 
 and weigh again. Connect 
 to the rest of the apparatus 
 and with careful heating (lamp) drive off as much of 
 the oxygen as possible; finally apply the full heat of a 
 Bunsen burner until no more oxygen comes away. Cool, 
 weigh tube and residue together, also measure the water 
 as in 1. 
 
 Fig. 12. 
 
78 
 
 EXERCISES IN CHEMISTRY. 
 
 Data and Computations. 
 Weight of tube + KClOa = 
 Weight of tube = 
 
 KCIO3 taken = 
 
 Weight of tube + KClOs = 
 Weight of tube + KCl =_ 
 Oxygen lost = 
 
 Percentage of O in KCIO3, by experiment : 
 
 Percentage of O in K.CIO3, computed from formula : 
 
 Error : 
 
 Number of c.c. of O obtained from gm. KClOs = 
 
 Number of c.c. of O corrected = 
 
 Number of c.c. of O * computed from " oxygen lost " = 
 Number of c.c. of O computed from formula KCIO3 = 
 * I litre of O weighs 1.43 gm. 
 
 3. The percentage of water of crystallization in certain 
 salts. 
 
 Apparatus : Two 20c.c. porcelain crucibles may be used 
 or two evaporating dishes of 30 to 50C.C. capacity. In this 
 and in several similar experiments it is most convenient and 
 accurate to weigh the dishes, and afterwards weigh the 
 results in the dishes, and subtract the 
 weights of the dishes. If the capacity 
 or delicacy of the balances do not per- 
 mit this, you can with sufficient care 
 scrape the substance to be weighed 
 from the dish with very little loss. Two 
 of the three following should be done. 
 A. Baric chlorid. Weigh one gramme 
 into one dish and two grammes into the other. Heat both, 
 at first very cautiously, till all action is over. Cool, — best 
 
 Fig. 13. 
 
EXPERIMENTS IN QUANTITATIVE RELATIONS. 79 
 
 in a dessicator (Fig. 13), so they will not absorb moisture, — 
 weigh : for certainty you should heat and weigh again. 
 
 Percentage of loss from i gm. = 
 
 Percentage of^ loss from 2 gm. =: 
 
 Percentage of water expressed in the formula, BaClj, SHjO =: 
 
 B. Copper sulfate. Like A. 
 
 Percentage of loss from i gm. = 
 
 Percentage of loss from 2 gm. = 
 
 Percentage of water expressed by CUSO4, SHjO = 
 
 C. Carbonate of Soda. Select perfectly fresh crystals 
 that have not efifioresced at all. Weigh rapidly. Heat with 
 greater caution than in A and B, since the crystals contain 
 so much water that they are likely to spatter. 
 
 Percentage of loss from i gm. = 
 
 Percentage of loss from 2 gm. =: 
 
 Percentage of water expressed by Na^CO:,, lOHjO = 
 
 4. T/ie quantity of salt in solution at a known tem- 
 perature. 
 
 Make a cold saturated solution of any of the following 
 salts : KCl, NaCl, KNO3, KCIO3, Pb(N03)2, HgCl^. Use 
 40 to 50C.C. of distilled water. You should make the solu- 
 tion the day before you wish to use it, and let it stand in a 
 covered beaker. Now weigh two dishes (crucibles, small 
 evaporating dishes or small beakers — see 3). Give your 
 solution a final stirring, take its temperature, and into each 
 dish measure with the utmost accuracy you can (with burette, 
 pipette, or graduate) (Fig. 14) loc.c. of solution; weigh 
 dishes and contents ; now on the water (Fig. 15), steam or 
 air bath, evaporate to dryness, stir and break up crystals 
 with glass rod, rinsing rod into dish with distilled water, if 
 necessary. A great difficulty with operations like this is 
 
8o 
 
 EXERCISES IN CHEMISTRY. 
 
 that the residue is not dry when it seems to be ; so let the 
 dishes stand in a dessicator, with concentrated HgSO^, over 
 
 Fig, 14. 
 
 night. Weigh dishes and contents. Now you have the 
 data to make your computations three ways. 
 
 in loo parts of solution by measure at 
 
 A. The weight of salt 
 temperature is 
 
 B. The weight of salt 
 perature is 
 
 C. The weight of salt 
 at temperature is 
 
 in loo parts of solution by weight at tem- 
 in loo parts of water by measure or weight 
 
EXPERIMENTS IN QUANTITATIVE RELATIONS. 8 1 
 
 The last is the more common method of treating solu- 
 tions. The first is also very important in normal solutions 
 used in analysis. 
 
 5. The solubility of a salt at different known temperatures. 
 
 The salts and methods are similar to 4. Weigh five small 
 dishes. Take the temperatures 20°, 30°, 40°, 50°, 6o°C. 
 
 200 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 LINE OF SOLUBILITY 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 190 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 1E0 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 170 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 100 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 ISO 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 140 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 130 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 ^110 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 Hlnn 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 Z 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 « y" 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 ^ JU. 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 70 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 50 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 _ 
 
 
 
 
 
 
 _ 
 
 _ 
 
 _ 
 
 _ 
 
 
 
 
 J 
 
 L 
 
 
 
 
 
 
 _ 
 
 
 
 _ 
 
 50 10° IS" 200 25° 300 35° 40° 45° SO" SS" 60° 65° 70° 75° 80» 85o 90° 'SBo 100° 
 TEMPERATURE 
 
 Make your saturated solution with about 6oc.c. of water at 
 io° above the highest temperature given, and with a ther- 
 mometer in the solution let it cool to 60°. Now take out, 
 
82 EXERCISES IN CHEMISTRY. 
 
 preferably with a pipette, loc.c. of the solution and put in 
 one dish ; weigh ; let the solution cool to 50°, and take 
 another loc.c., etc. Of course you may select other tem- 
 peratures, but you will find it difficult to manage the solution 
 near the boiling point. Rinse your pipette with boiling 
 water before and after using. Dry your dishes as in 4. 
 Make your calculations as in 4, C. Then lay off the solu- 
 bility in 100 parts of water at each temperature upon 
 squared paper, and draw the curve or straight line. 
 
 In 100 of water, salt at 20° = 
 
 ^t 30° = 
 at 40° = 
 at 50° = 
 at 6a° = 
 
 6. The proportion of oxygen in air, by volume. 
 
 Take a soft glass tube \ or | in. in diameter, 18 to 20 in. 
 long ; bend it like Fig. 6. (The upper end is drawn out, 
 closed, and bent over about i in. ; the lower end is an open 
 loop about 2 in. long.) When the tube is cool, slip to the 
 upper bend a piece of phosphorus half the size of a small 
 pea. (Remember the inflammability, etc., of phosphorus.) 
 Stand the tube in a jar of water so that the loop will fill 
 with water ; mark the height of the water in the long arm 
 with a file or string. Return the tube to the jar, fill the jar 
 up with water. Now gently heat the phosphorus with the 
 lamp flame for about five minutes, taking care not to drive 
 any air around the loop ; let the tube stand in the water till it 
 has cooled to the temperature of the room. Make the level 
 in the long arm coincide with level in water jar ; mark the 
 height of the water. Now measure (i) contents of tube 
 before heating ; (2) contents of tube after heating. Reduce 
 
EXPERIMENTS IN QUANTITATIVE RELATIONS. 83 
 
 both measurements to the same temperature, if the temper- 
 ature of the room has changed sensibly during the experi- 
 ment. 
 
 Air taken = 
 
 Residue ^ 
 
 Percentage lost ( Oxygen) =: 
 
 1. Tke behavior of indicators {preliminary to 8). 
 
 Apparatus : Perfectly clean test tubes, distilled water, 
 all the laboratory indicators — litmus, cochineal, lakmoid, 
 phenol-phthalein, etc. In one test tube very dilute NaOH, 
 in another very dilute HCl. 
 
 Now in as many test tubes as you have indicators put 
 5c. c. of water and add two or three drops of the various 
 indicators. Note colors (neutral). Add one drop of HCl to 
 each ; shake thoroughly ; note colors (acid) ; add, drop by 
 drop, NaOH to each, shaking after each drop ; note colors 
 (alkaline) ; once more add HCl to note the change from 
 alkaline to acid. (These changes are best obtained by 
 dropping from a burette.) 
 
 Sharp 
 Number. Name Neutral. Acid. Alkaline. or Indistinct. 
 
 8. A. The relative neutralizing power of a certain 
 quantity, and twice that quantity of an alkali (Na^CO^, lO/f^O) 
 with the acids H^SO^ and HCl. 
 
 Weigh and dissolve in soc.c. of water 2 gm. and 4 gm. of 
 NagCOg, IOH2O, using only perfectly fresh crystals. Add 
 to each about 5 drops of cochineal solution (use as little as 
 will give a perceptible color). In the most accurate meas- 
 uring apparatus at hand — burette, pipette or cylinder — take 
 
84 EXERCISES IN CHEMISTRY. 
 
 a measured quantity of quite dilute HgSO^ (the ordinary 
 dilute acid of the laboratory plus 3 or 4 parts of water). 
 Now neutralize the two quantities of carbonate of soda, 
 with constant stirring, and note the quantity of acid required' 
 for each. 
 
 2 gm. NajCOa require c.c. dilute H2SO4. 
 
 4 gm. NagCOa require c.c. dilute H2SO4. 
 
 8. B. Exactly the same, using quite dilute HCl instead 
 
 of H2SO4. 
 
 2 gm. NagCOa require c.c. dilute HCl. 
 
 4 gm. NajCOa require c.c. dilute HCl. 
 
 9. The relative neutralizing powers of sodium kydroxid 
 and potassium kydroxid, with HCl 
 
 Weigh as rapidly as possible, in weighed beakers or watch 
 glasses, about J in. from fresh sticks of the best NaOH and 
 KOH obtainable. Dissolve in water and make each up to 
 a convenient quantity — say looc.c. Now take of the 
 NaOH solution enough to contain exactly .4 gm. of NaOH 
 (quantity weighed : .4^ lOO : x. x = no. of c.c). Color 
 with an indicator. (In this particular case you will find 
 phenol-phthalein very satisfactory ; litmus, cochineal, lak- 
 moid may be used.) Put in your measuring apparatus a 
 definite quantity of dilute HCl (as in 8) and find how many 
 c.c. are required to neutralize exactly the NaOH. Number 
 of c.c. dilute HCl for .4 gm. NaOH := . Now 
 
 measure exactly this number of c.c. of HCl into a beaker; 
 rinse burette, first with water, then with a little of the 
 KOH solution, and fill with the latter. Now neutralize the 
 HCl with the KOH. No. of c.c. KOH = 
 
 Calculate what weight of KOH was used (No. of c.c. X 
 KOH weighed 
 
 100 ■'■ 
 
EXPERIMENTS IN QUANTITATIVE RELATIONS. 
 
 85 
 
 The quantity of HCl which neutralizes .4 gm. NaOH, 
 neutralizes KOH. 
 
 * In what relation do these numbers stand to the formula 
 weights of NaOH and KOH ? 
 
 10. A. The quantity of sulfate, chlorid and nitrate of 
 soda from a given weight of carbonate ; B. also by calculation 
 the quantities of the acids required ; C. and the relations of 
 these quantities to the formula weights. 
 
 A. I. Weigh 2.86 gm. of NaaCOg, lOH^O ; dissolve in 
 20 to 30C.C. of distilled water in a small (weighed) evaporat- 
 ing dish; color with 
 3 or 4 drops of lit- 
 mus; now add, drop 
 by drop, dilute 
 H2SO4 until the 
 liquid is pink after 
 stirring. Evapo- 
 rate to dryness 
 upon a water (Fig. 
 15) or hot-air bath. 
 Whenever during 
 the evaporation the =^==rg 
 blue color returns 
 to any part of the 
 liquid, add just 
 enough acid to re- 
 store the pink. Do 
 not add any excess of acid 
 
 * Note. — If the manipulation has been careful, the results will depend upon 
 the purity of the hydroxids. In this laboratory the NaOH of the best makers 
 has been found nearly pure; KOH commonly corresponds very closely to the 
 formula KOH, H^O. 
 
 pnn»r 
 
 Fig. 15. 
 
 After the contents of the dish 
 
86 EXERCISES IN CHEMISTRY. 
 
 seem to be dry, continue the heat for an hour. Finally 
 heat a minute or two with a Bunsen. (If the contents of 
 the dish do not become dry on the water bath, it is because 
 too much H2SO4 has been added ; excess can be driven 
 out by heating in a hood, but there is danger of loss by 
 spattering.) Cool ; weigh. 
 
 Weight of Na^SO, obtained from 2.86 gm. NaaCOa, lOH^O = 
 Weight of Na2S04 theoretically obtainable = 
 
 B. Na^SOi obtained : N^^SOi used (tC) = formula weight Na2S04 : formula 
 weight HaSO,. 
 
 Required to neutralize 2.86 gm. NajCOa, IOH2O gm. HjSOi. 
 
 A. 2. Repeat this experiment with this difference : Use 
 no indicator, and add to the solution of NajCOg, lOHgO, 
 5C.C. of concentrated pure HCI, very slowly. 
 
 Weight of NaCl obtained from 2.86 gm. NaaCOa, lOHjO = 
 Weight of NaCl theoretically obtainable = 
 
 B. NaCl obtained : ^C/«j^(/(x)=: formula weight NaCl: formula weight HCI. 
 Required to neutralize 2.86 gm. NajCOs, lOHjO gm. HCI. 
 
 A. 3. Repeat this as in A, 2 with 2c.c. concentrated pure 
 HNO3. 
 
 Weight of NaNOs obtained from 2.86 gm. NajCOa, lOHjO = 
 Weight of NaNOs theoretically obtainable = 
 
 B. NaNOs obtained: HNO's used (x) = formula weight NaNOs : formula 
 weight HNO3. 
 
 Tabulating: 2.86 gm. NajCOj, IOH2O gave gm. Na^SOj. 
 
 gm. NaCl. 
 gm. NaNOj. 
 2.86 gm. NajCOa, lOHjO required gm. H2SO4. 
 
 Formula weight of H2SO4 = 
 
 HC1 = 
 
 HNO» = 
 
 gm. HCI. 
 gm. HNO3. 
 
 C. How does the neutralizing power of the formula weight of H2SO4 com- 
 pare with that of the other aciils, and hpw do they compare with each other? 
 
EXPERIMENTS IN QUANTITATIVE RELATIONS. 87 
 
 11. A. The weight of copper precipitated from solution 
 by a definite weight of another metal {cadmium). 
 
 Hold a rod of cadmium in the. Bunsen flame and as it 
 melts let the drops fall from a height of 6 or 8 in. upon a 
 smooth surface, like stone or slate, so that they will spread 
 out into thin disks. 
 
 Dissolve s gm. of CUSO4, SH^O in soc.c. of water; add 
 2 drops of dilute. H2SO4. Weigh exactly 1.12 gm. of the 
 thin cadmium, put it in the solution and let it stand 24 
 hours, or until every particle of the cadmium is dissolved. 
 Now wash the precipitated copper by decantation four times, 
 taking care to lose no Cu, and also to pour off the water 
 completely each time. Dry the Cu (water bath or by 
 letting it stand in dessicator), weigh, either in the beaker, 
 or by brushing out every particle. 
 
 Cadmium taken I.I 2 gm. Cu obtained gm. 
 
 Look up equivalents and atomic weights of Cd and Cu. 
 
 11. B. The weight of lead precipitated by cadm.ium. 
 
 The experiment is similar to 11, A. Use 6 gm. of lead 
 acetate, add 2 or 3 drops of acetic acid — enough to make 
 the solution clear. Only the thinnest portion of the Cd 
 will do for this, and before washing, you must make sure that 
 no undissolved Cd is left. 
 
 Cd taken 1. 1 2 gm. Pb obtained gm. 
 
 C. and D. The same method may be applied to AgNOg 
 and HgClg. There is, however, more danger of loss of fine 
 particles, and the weighing must be done in the beaker. 
 
EXERCISES IN CHEMISTRY. 
 
 [13. Hydrochloric acid gas is one-half hydrogen {and one- 
 half chloriti) by measure. 
 
 Apparatus. See Figures. A thick glass 500C.C. Erlen- 
 meyer filtering flask with side neck, provided with tight rub- 
 ber stopper, about i ft. 
 of rubber tubing for 
 the side neck and a stout 
 pinchcock, which must 
 close the tubing air- 
 tight. A jar holding 300 
 to 400c. c. of water. Ap- 
 paratus for making HCl 
 as below. (Fig. 16.) 
 
 Everything must be 
 perfectly dry and the 
 joints and stoppers must 
 be air-tight. 
 
 Make HCl by dropping 
 concentrated HjSO^ 
 from a stoppered funnel 
 upon rock salt in an ordinary or Erlenmeyer flask. Occa- 
 sionally warm the HCl flask ; now fill the side-necked 
 flask with HCl gas by downward displacement, continuing 
 a steady current not less than half an hour (hood). Drop 
 3C.C. of Mg powder (photographic) into the side-necked 
 flask, close tight, put the free end of the rubber tube 
 in the jar of water. Now very cautiously let water suck 
 into the flask, to about half its capacity, by opening the 
 pinchcock. (If the rubber tube was full of HCl, the action 
 will be very lively, otherwise it will be slow, until the 
 first drop of water gets over.) Close the pinchcock tight 
 and let the flask stand until action ceases (12 hours or more). 
 
 Fig. 16. 
 
EXPERIMENTS IN QUANTITATIVE RELATIONS. 
 
 89 
 
 Finally tip the flask over so that the water will come into 
 the neck (Fig. 17), open the pinchcock and let the water 
 in the flask and jar come to equal pressure. When this 
 condition is reached, set the flask upright again. Measure 
 the water in it, and its total contents, making the necessary 
 
 Fig. 17. 
 
 corrections. The flask was full of HCl, the HCl was 
 absorbed by water and afterwards decomposed by Mg free- 
 ing the H. The dry gas has almost no action upon Mg. 
 
 This experiment needs very intelligent and skilful manip- 
 ulation.] 
 
 Contents of flask, HCl 
 Water, representing CI 
 Gas H 
 
 c.c. 
 
 C.C. 
 
 c.c. 
 
go EXERCISES IN CHEMISTRY. 
 
 13. The relative volumes of CO and CO^ obtained by the 
 decomposition of ferrous oxalate. 
 
 Apparatus : Exactly that used in Experiment 3. Weigh 
 just 1.5 gm. of well-dried FeC204 into the ignition tube; 
 apply gradually the full heat of a Bunsen burner ; continue 
 to heat until the residue in the tube is of uniform brownish 
 black and all gas has ceased to come away. After cooling 
 measure the water driven over into the jar, with tbe pre- 
 cautions and corrections previously described. 
 
 Now this experiment is to be repeated exactly, with one 
 important difference in the conditions, viz. before beginning, 
 dissolve 5 to 10 gm. of NaOH (or KOH) in the water of 
 the flask ; during ignition and cooling shake round the water 
 in the flask. 
 
 Total volume of gas (CO + CO2) = c.c. 
 
 Volume not absorbed by NaOH (CO) = c.c. 
 
 Volume absorbed by NaOH (CO2) = c.c. 
 
 It appears that in the decomposition of FeC204, CO and 
 CO2 are produced in volumes. 
 
 Now, assuming that the equation FeCgO^ =: FeO + CO + 
 CO2 is correct, which occupies the most space — 28 parts (by 
 weight) of CO or 44 of CO2 •■' Which occupies the most 
 space — a certain number of molecules of CO, or the same 
 number of molecules of CO2 .' 
 
 14. The weight relations of certain com.pounds {of lead) 
 obtained by successive transformations. 
 
 Try to account for every slight variation from the result 
 expected. 
 
 A. The buff oxid {PbO) from white lead l^PbCO^, 
 PbO^H^. Weigh just 10 gm. of the best, dry, white lead 
 
EXPERIMENTS IN QUANTITATIVE REI-ATIONS. 9I 
 
 obtainable into an evaporating dish. Apply small heat and 
 stir gently ; after a time increase the heat ; continue to heat 
 and stir until the product is a powder of uniform buff color, 
 showing no streaks of white when the little lumps are 
 smoothed out with the rod. Cool ; weigh. 
 
 10 gm. white lead give gm. PbO. 
 
 If pure 2PbC03, PbOjHj theoretical yield is gm. PbO. 
 
 Difference := 
 
 B. The yellow oxid (PbO) from the buff oxid {PbO). In 
 the same dish continue to heat the product just weighed; 
 stir, loosen from the bottom, stop heating when the contents 
 of the dish arc of uniform bright yellow ; cool ; weigh. 
 
 gm. buff PbO give gm. yellow PbO. 
 
 Theoretical yield gm. yellow PbO. 
 
 Difference = 
 
 C. The nitrate {Pb{NO^^ from the yellow oxid {PbO). 
 To the weighed PbO in the dish add with stirring and 
 gentle warming HNO3 (i part HNO3, i part water), until 
 the yellow powder is either dissolved or changed to a white 
 compound. Now put the dish in water or air bath and 
 bring to dryness ; break up the crystals ; dry thoroughly in 
 a dessicator ; weigh. 
 
 gm. PbO give gm. Pb(N03)j. 
 
 Theoretical yield ^gm. Pb(N03)j. 
 
 Difference = 
 
 D. The sulfate (PbSO^) from the nitrate {Pb{NO^^. 
 Dissolve the weighed Pb(N03)2 in as little cold water as 
 convenient (200 c.c. or so in a beaker) ; add, with constant 
 stirring, dilute H2SO4, until an additional drop gives no 
 further precipitate (let it settle before testing) ; now let it 
 stand until perfectly clear above the white precipitate ; pour 
 
92 EXERCISES IN CHEMISTRY. 
 
 off the liquid, steadily but rapidly, down to the edge of the 
 precipitate ; fill with water again and allow to settle; decant; 
 repeat; put the beaker where the precipitate will be thor- 
 oughly dried* ; weigh. 
 
 gm. PbCNOOa give gm. PbSOj. 
 
 Theoretical yield gm. PbSOj. 
 
 Difference = 
 
 lo gm. of white lead have given gm. PbS04. 
 
 Theoretical yield, direct, gm. PbS04- 
 
 Difference = 
 
 * Note. — A little more accurate results could be obtained by adding one-half 
 its volume of alcohol to the water and also by decanting through a filter and 
 brushing off the precipitate when dry. If the precipitate does not dry well, it 
 may be moistened with alcohol and dried again. 
 
 SUMMARY. 
 
 After completing as many of the experiments in Part III 
 as seem proper to the instructor, the student is expected to 
 prepare a complete tabulated summary and to follow this 
 with the statement of what fact or principle he has learned 
 from each experiment, with the reasons why absolute accu- 
 racy was not to be expected, and with any suggestions he can 
 make for greater accuracy under the conditions of class 
 work. 
 
PART IV. 
 
 SUGGESTIONS UPON PREPARATIONS. 
 
 References. 
 
 1. A solution of approximate strength. 
 
 2. A solution of exact strength. 
 
 3. A saturated solution. 
 
 4. A solution of a liquid. 
 
 5. A solution of definite specific gravity. 
 
 6. The distillation of liquids. 
 
 7. The preparation and purification of gases. 
 
 8. The preparation of a solid by precipitation. 
 
 9. Crystalline salts. 
 10. Sublimation. 
 
 SUGGESTIONS UPOK CHEMICAL PREPAMATIONS. 
 
 The preparation of reagents for laboratory use, chiefly in 
 the form of water solutions, is of daily necessity and may 
 often be made a profitable exercise for the beginner. More- 
 over the manufacture of many chemicals, especially of 
 crystalline salts, is exceedingly interesting and occasionally 
 very convenient ; though it is seldom economical if the 
 article can be purchased. 
 
 The field is very broad and only a few instructions 
 applicable to the simpler processes can be given here. 
 Among books that will be found useful in this field are 
 Erdmann's " Chemischer Praparate," Curtman's " Chemical 
 Reagents," Attfield's Chemistry, many of the larger text- 
 
94 EXERCISES IN CHEMISTRY. 
 
 books, and also the manuals of Qualitative and Quantitative 
 Analysis. 
 
 In this book opportunity for selection is found in the 
 list of chemicals, Part I ; the residues of many experi- 
 ments collected from a number of pupils : viz. Part I, 7, i, 
 8, I, 11, 4. 13, I, 14, I, 18, I, 19, I, 24, 2, 26, i, 
 34, 4, 35, 2, 36, 3, 39, i ; Part III, 1, 4, 5, 8, 9, 10, 
 13, 14 ; larger quantities in many of the experiments given 
 or referred to in Part I, 43-49. 
 
 The only additional apparatus absolutely needed is 
 evaporating dishes, holding i to 3 litres, funnels, about 6 in. 
 filters, 10 to 12 in., thin glass crystallizing dishes 3 to 6 in. 
 
 1. The preparation of a solution of approximate strength, 
 of a salt in water. Weigh the salt in grammes, measure 
 the water in centimeters ; mix ; the solution may be 
 hastened by stirring, pulverizing the salt, or by gently 
 warming the water ; whenever there is a residue of dirt or 
 foreign material the solution is to be filtered before using. 
 When it is destructive to filter paper, it may be decanted, 
 syphoned off, or passed through a funnel plugged with glass 
 wool. 
 
 When a salt is merely to be dissolved in water without 
 specification of quantity, the novice is likely to use too little 
 water ; ten to twenty parts of water should generally be 
 used to one of the salt. 
 
 3. The preparation of a solution of exact strength. This 
 can only be done directly when the substance can be 
 obtained pure and can be weighed exactly. 
 
 Weigh the substance and dissolve it in a measuring flask 
 with a little water ; after solution add water enough to make 
 an exact measure of the whole solution. Thus -^^ normal 
 AgNOg solution may be made by dissolving 17 gm. of 
 
SUGGESTIONS UPON PREPARATIONS. 
 
 95 
 
 AgNOg in water in a litre flask and then filling the flask 
 exactly to the mark. 
 
 (In the case of substances like liquid acids and the 
 alkalis, not exactly weighable, solutions are made, by esti- 
 mate, a little over the desired strength ; their exact strength 
 is determined by methods given in the manuals on Volu- 
 metric Analysis, then they are (or may be) diluted with 
 water, by computation, to the exact strength desired.) 
 
 3. The preparation of a saturated solution. Merely take 
 care that undissolved salt remains constantly at the bottom 
 of the bottle. The solution will take care of itself, varying 
 its strength with the variations in the temperature of the 
 room. In the case of lime water and some very sparingly 
 soluble substances the first water is poured off and thrown 
 away. 
 
 4. A solution of a liquid. A measured volume of the 
 liquid is poured into as many measured volumes of water 
 as will give the required strength, thus : to make dilute 
 HjSO^ add with great caution i part of 
 concentrated HjSO^ to 4 parts of water. 
 
 5 . A solution of definite specific gravity. 
 This requires a float or hydrometer mark- 
 ing specific gravity (Fig. 18). (It is also 
 done with a Westphal balance.) Thus to 
 make HNO3 specific gravity 1.2 put the 
 hydrometer in a jar tall enough to float it, 
 add HNO3 till it floats, and then water 
 till its water line is 1.2. The water should 
 be added slowly and with stirring. 
 
 In the case of salts make a solution over the required 
 strength and dilute. 
 
 6. The preparation or purification of a liquid by distilla- 
 
 Fig. 18. 
 
96 
 
 EXERCISES IN CHEMISTRY. 
 
 Hon. (See Fig. 9, also "Distillation " in larger books.) The 
 liquid or the materials to make it are put into a side-necked 
 round-bottomed flask, never making it more than one-half 
 full, a thermometer passes through the cork of the flask so 
 that its bulb is opposite the side-neck, the side-neck passes 
 through a little cork into the inner tube of a condensing 
 
 Fig. 19. 
 
 apparatus. Heat is applied through a sand bath or wire 
 gauze, just enough to keep the liquid gently boiling ; water 
 is passed through the condenser (Fig. ig), and the distillate 
 is collected between certain definite temperatures on the 
 thermometer, depending on the substance. In the case of 
 corrosive liquids, like nitric acid, a glass-stoppered flask or 
 retort must be used. As most liquids subject to distillation 
 
SUGGESTIONS UPON PREPARATIONS. 
 
 97 
 
 are highly inflammable this should not be undertaken with- 
 out special instructions. There are many forms of distilling 
 apparatus. 
 
 7. The preparation and purification of gases. Sufficient 
 for the present purpose has been given upon preparation by 
 the methods employed in Part I. (Let the pupils look over 
 
 Fig. 20. 
 
 Parts I and III and describe Preparation and Collection of 
 Gases.) The side-necked test tube is, with care, very satis- 
 factory for small quantities of gases made either from solids 
 or liquids ; as larger quantities are required, flasks with 
 safety tubes are much better ; the Erlenmeyer flask is con: 
 venient because of its broad base. Gases soluble in water 
 are collected by upward or downward displacement of air, 
 over brine, or over mercury. It is often necessary to purify 
 a gas, and this is done by causing- it to pass through some 
 substance which absorbs the impurities (Fig. 20). For in- 
 stance, the gas is often washed simply by passing it through 
 water ; vapor of water is removed — the gas is dried — by 
 concentrated H2SO4. A short delivery tube from the gas 
 
gS EXERCISES IN CHEMISTRY. 
 
 generator dips under the liquid in another vessel, and a long 
 delivery runs from this second vessel, which may be a side- 
 necked test tube, or flask, or Wolff bottle. When the gas 
 is to be purified by a solid the latter, in a loose and granular 
 condition, is placed in a U-shaped tube, or in a " chlorid of 
 calcium tube." 
 
 8- The preparation of an insoluble compound by precipi- 
 tation. (See Part I, 49.) The two substances which react 
 to produce the precipitate are in cold water solution gen- 
 erally. One is added to the other with constant stirring 
 until it produces no more effect. To know that enough has 
 been added either let the mixture settle until it is clear, or 
 filter a few drops, and add a drop more of the reagent. It 
 is often convenient to make an approximate computation 
 of the quantities necessary. The liquid contains the other 
 product of the reaction, so it is necessary to free the pre- 
 cipitate from this by washing. This may be done by decan- 
 tation — pouring off the clear liquid down close to the pre- 
 cipitate and refilling with water at least five times ; or by 
 bringing the precipitate upon a filter and washing, letting 
 each portion of water run through before adding another. 
 Precipitates settle and filter better after thorough stirring. 
 They may be boiled and the wash water may be hot when 
 there is not much difference between their solubility in cold 
 and hot water. To know the precipitate is clean the last 
 drops of wash water are tested with some reagent for the 
 substance being washed out ; e. g. if it is a chlorid, with 
 AgNOj, if it is a sulfate, with BaClg. if it is an acid or alkali, 
 with litmus. The next step is to dry the precipitate ; this 
 is done by keeping it in a warm place — register, steam 
 oven, evaporating dish over water bath — and protecting it 
 from dust. This is a slow process. Finally it is gently 
 
SUGGESTIONS UPON PREPARATIONS. 99 
 
 separated from the filter paper. To illustrate. Prepare 
 about 10 gm. BaCOg. 
 
 BaClg, 2H2O + Na^COg, lOH^O = BaCOg + 2NaCl + I2H2O. 
 
 244 286 197 
 
 We see that 10 gm. BaCOg will require about 13 gm. of 
 the chlorid, and that 15 gm. of the crystallized carbonate 
 will be more than enough ; dissolve them separately in cold 
 water ; proceed as above ; common salt (NaCl) is the product 
 to be washed out and BaCOg is insoluble in hot water ; wash 
 therefore with hot water until a few drops tested with 
 AgNOg give only the faintest cloud. 
 
 Again, make some fresh lead sulfid. Pb(N03)2 +. HgS = 
 PbS+2HNOg. Dissolve a gram or so of Pb(N03)2 in 
 cold water ; pass H^S through the solution till the precipi- 
 tate seems ready to settle in flakes and leave the liquid 
 clear; filter and wash until the wash water is not acid to 
 litmus. For use in 11, 4, Part I,_it is not desirable to have 
 it dry. 
 
 9. Tke preparation of crystalline salts. Among the 
 general methods of preparation are the action of acids upon 
 metals, metallic oxids, hydroxids, and carbonates ; the action 
 of sulfuric acid upon the salts of many other acids ; nearly 
 exact double decomposition from which one of the products 
 is insoluble and the other is soluble ; direct union of chlorin 
 with metals ; the action of dry heat upon salts which give 
 up part or the whole of their oxygen. 
 
 Each salt has individual characteristics which can only be 
 learned by careful study. 
 
 In general crystalline salts are obtained by evaporating 
 the water from their solutions. 
 
 If the crystals form too rapidly and imperfectly, the 
 
lOO EXERCISES IN CHEMISTRY. 
 
 evaporation has gone too far ; they should be redissolved in 
 a little more hot water. 
 
 If they delay too long in forming, the evaporation has 
 not gone far enough ; this can be remedied by further 
 evaporation, or by waiting for the water to evaporate spon- 
 taneously. 
 
 Salts which are very much more soluble in hot than in 
 cold water will often crystallize rapidly from boiling water 
 as soon as the heat is removed. 
 
 When the size of the crystals is of little importance the 
 solution may be evaporated until a slight crust begins to 
 show upon the sur- 
 face of the liquid ; f--^ 
 
 then the heat is 
 removed. 
 
 Or a drop may be 
 put on a cool glass 
 surface from time to time ; when it quickly shows minute 
 crystals the heat should be removed. 
 
 Spontaneous evaporation in a warm room, or with heat 
 insufficient to cause the liquid to boil, often results in large 
 crystals. 
 
 Evaporation to dryness is to be avoided ; it may cause 
 spattering, it spoils the crystals, it prevents sometimes 
 their combination with water, it sometimes melts the salts 
 and it leaves no, mother liquor containing impurities to 
 pour off. 
 
 Stirring or agitation of a liquid at the point of crystalliza- 
 tion usually hastens the formation of small crystals. Some- 
 times this method is used to secure a crop of uniform size. 
 
 The solution at the point of crystallization is left in the 
 evaporating dish to cool and evaporate spontaneously, or it 
 
 Fig. 21. 
 
SUGGESTIONS UPON PREPARATIONS. lOI 
 
 is poured into a crystallizing dish (Fig. 21) which has just 
 been rinsed with boiling water. 
 
 More perfect forms may often be obtained by turning over 
 some of the better crystals from time to time. 
 
 And in some cases very perfect crystals may be made by 
 pouring the solution into a tall narrow jar and letting a 
 weighted string run from the top to the bottom of the jar, 
 not touching the sides. 
 
 On the manufacturing scale, or when the same process is 
 to be often repeated, it is convenient to evaporate to some 
 definite specific gravity, ascertained by a float or hydro- 
 meter. 
 
 The evaporation of small quantities of water is nicely 
 effected under a bell jar in the presence of a beaker of 
 concentrated HjSO^. 
 
 After a sufficient crop has been formed the mother liquor 
 is poured off, the crystals are loosened with a glass rod 
 and transferred to a funnel, or to a porcelain colander, and 
 washed with a very little cold water and allowed to drain. 
 The mother liquor and wash water may be evaporated for 
 more crystals. Of course very soluble salts will not bear 
 much washing. 
 
 After the crystals have drained well and seem fairly dry 
 they should, if efflorescent, be bottled. Ordinary drying 
 may be effected by spreading the crystals upon a glass plate 
 or a plate of biscuit ware. Thorough drying even of many 
 deliquescent salts may be effected by. the bell jar and con- 
 centrated H2SO4. 
 
 Recrystallization consists in dissolving the crop obtained in 
 a convenient quantity of boiling water and evaporating again. 
 Among the salts, products of Part I, to which these prin- 
 ciples apply rather simply, are : ZnS04, THjO ; CUSO4, 
 
102 EXERCISES IN CHEMISTRY. 
 
 SH^O; FeSO^, TH^O ; Na^SO^; KCl, NaCl, (NH4)2S04, 
 NH4NO3, Pb(N03)2. 
 
 Many double salts crystallize beautifully ; e. g. (N 114)2504, 
 FeS04, 6H2O. Dissolve with heat 4 gm. of iron (4 three- 
 penny nails) in H2SO4 (4OC.C. of water, 8 to loc.c, and no 
 more, of concentrated H2SO4). After the solution is com- 
 plete, add dilute NH4OH to the liquid until a slight precipi- 
 tate remains after stirring. Add dilute H2SO4, drop by 
 drop, until the liquid is again acid. Filter into an evapo- 
 rating dish and concentrate until a drop gives abundant 
 crystals as it cools on a glass surface. Now transfer the hot 
 liquid to a crystallizing dish or wide beaker, and let it crys- 
 tallize. Drain off the mother liquor, break up the mass of 
 crystals, and wash them twice with a very little cold water, 
 then dry them. 
 
 After a very little practice the pupil will find excellent 
 opportunity for original inquiry in seeking just the" right 
 conditions with any particular salt. 
 
 10. Purification or crystallization by sublimation is applic- 
 able to solids that can be changed to vapor by heat. 
 
 The salt is put into a dry evaporating dish without lip, 
 
 or in a watch glass ; a filter paper is put over the dish and 
 
 an inverted funnel, or watch glass, placed 
 
 above (Fig. 22). The lower dish is gently 
 
 Fig. 22. heated, — best in a sand bath, — the solid 
 
 changes to vapor, passes through the paper and condenses. 
 
 This is applicable to small quantities only. 
 
 Illustration : — Hgig, by precipitation and sublimation. 
 Dissolve a drop of Hg in 2 or 30.0. of nitro-hydrochloric 
 acid (i part HNO3, 3 parts HCl) with heat. After all has 
 dissolved that will, dilute with 10 to 20C.C. of water, and 
 filter into a beaker. (Solution of HgCl2 or Hg(N03)2 may 
 
SUGGESTIONS UPON PREPARATIONS. IO3 
 
 be used instead.) Now add with care a solution of KI. 
 Precipitate is Hglj, easily soluble in excess of KI. Filter, 
 wash precipitate four or five times with very small quantities 
 of cold water, letting each portion run through completely. 
 Set precipitate aside to dry in the filter. Remove the dry 
 precipitate to a watch glass, cover it with another, place on 
 iron, and very gently heat. Notice the color changes. The 
 Hglj sublimes to the upper glass. Yellow crystals are 
 rhombic ; red, octahedral. 
 
PART V. 
 
 CHEMICAL TERMS. 
 
 The student of chemistry must acquire early familiarity 
 with the use of the following terms, in order to read intelli- 
 gently and to speak accurately. 
 
 In many cases it must be understood that the definition 
 is suggestive rather than complete. , 
 
 A — in composition; not; without; e. g. amorphous, not crystalline; anhy- 
 drous, without water. 
 
 Acid. A substance which neutralizes a base; gour; turns litmus red. 
 
 Acid Salt. A salt in which there remains displaceable hydrogen; e. g. 
 HNaSOj. 
 
 Alkali. Often applied to a base that dissolves in water. 
 
 Alkaline. Like an alkali; turns litmus blue. 
 
 Allotropy. The property of an element to exist in forms with different 
 characteristics. 
 
 Alloy. Physical or chemical union of metals. 
 
 Amalgam. Alloy in which mercury is one of the metals. 
 
 Amorphous. Not crystalline. 
 
 Analysis. Separation into parts; applied also to the determination of the 
 presence or quantity of substances by chemical methods. 
 
 Anhydride, An oxid capable of changing to an acid (or alkali) by the addi- 
 tion of water. SO2 -f- H2O = H2SO3. Sulfurous anhydride and water give 
 sulfurous acid. 
 
 Anhydrous. Without water; applied to salts that have been deprived of 
 water of crystallization ; also to water-free substances in general. 
 
 — ate. Termination of ternary compound ; implies more oxygen than — ite. 
 
 Atom. The unit of chemical change, differing for each element. 
 
 Atomic Weight (or Mass). The relative weight in which, or in multiples of 
 which, an element always takes part in chemical changes. 
 
 Base. A substance which neutralizes an acid; usually a metal or the oxid, 
 hydroxid, or carbonate of a metal. 
 
 Basicity. The number of parts (or atoms) of hydrogen in an acid which can 
 be displaced by a metal. The basicity of HCl is i, of H2SO4 is 2. 
 
 Basic Salt. A salt having basic properties; capable of neutralizing acid. 
 
 Z04 
 
CHEMICAL TERMS. IO5 
 
 Bi — Prefix meaning two, twice, double. 
 
 Binary. Compound consisting of two elements. 
 
 Boil. Changing to vapor with visible escape of bubbles. 
 
 Bond. The valence of an element represented by a line. H — , — O — . 
 
 Centigrade Scale. Freezing point .at 0°, boiling point at 100°. 
 C=|(F — 32), F=fC + 32. 
 
 Centimeter, cubic. The bulk of i gram of water; j-p^jj litre; ^ teaspoonful, 
 ■^f fluid oz., approximately. 
 
 Chemical Attraction. (Chemism, Chemical Affinity.) The energy which 
 causes and maintains chemical union. 
 
 Combustible. Capable of burning, usually with oxygen. 
 
 Combustion. Rapid chemical union with light and heat; as usually understood, 
 rapid union with oxygen. 
 
 Commercial. - As commonly sold ; without elaborate purification. 
 
 Compound. The product of the union of any two or more elements; a decom- 
 posable substance. 
 
 Concentrated. Strong; in full strength; not weakened by dilution. 
 
 C. P. Chemically pure ; carefully freed from foreign substances. 
 
 Crith. The weight of 1 litre ot hydrogen — .0896 gm. • 
 
 Crystalline. Bounded by geometrical surfaces; more or less regular and 
 glistening. 
 
 Decant. To separate a liquid from a solid by careful pouring. 
 
 Decompose. To separate into simpler parts. 
 
 Decrepitate. Snapping to pieces when heated. 
 
 Deflagration. Burning with flashes of light, like moist gunpowder. 
 
 Deliquescent. Becoming liquid; applied to substances that gather moisture 
 from the air. 
 
 Di — prefix, like bi — . 
 
 Diffusion. The intermingling of gases. 
 
 Dilute. Weakened by the addition of water. 
 
 Displacement. Collection of a gas by sending it downwards or upwards into 
 a jar of air. 
 
 Dissociation. Breaking into simpler form and occupying more space. 
 
 Distillation. Changing a liquid to vapor by heat, condensing and collecting 
 it again. 
 
 Double decomposition. The exchange of like parts between two com- 
 pounds. See p. 54. 
 
 Double Salt. A salt containing two metals. 
 
 Dyad. Element with valence 2. 
 
 Effervescence. Escape of a gas in bubbles, usually from chemical change. 
 
I06 EXERCISES IN CHEMISTRY. 
 
 Efflorescent. Becoming a powder; applied to salts that lose their water of 
 
 crystallization at ordinary temperature. 
 Element. One of the 72 or 73 forms of matter at present not separable into 
 
 simpler parts. 
 Empirical. Experimental. 
 Equation. See pp. 14, 65, 68. 
 Equivalent. Adject., having equal chemical effect in neutralizing, precipitating, 
 
 etc.; noun, the weight defined in Part II, 4, B & C. 
 Evaporate. To change to vapor; to boil away. 
 Excess. " In excess," adding a reagent in sufficient quantity to have a little left 
 
 over; as in adding an acid to an alkali, add till the whole liquid is acid. 
 Filtrate. The liquid which runs through the filter. 
 Formula. The expression of a compound and its relative composition. 
 Gram. Weight of cubic centimeter of water; ■j-j'gTf kilogram; 1000 milligrams; 
 
 15-434 grains. 
 
 Graphical. Representing structure by diagrams | H I — | O | — | H | 
 
 Gravimetric. By weight. 
 
 Halogens. The chlorin group of elements. 
 
 Hard.- Applied to water containing salts of lime or magnesium. 
 
 Hydrate. Hydroxid; better usage, a compound with H2O. 
 
 Hydrated. Combined with water. 
 
 Hydro — Prefix applied to acids containing no oxygen; HCl , hydrochloric acid; 
 
 HCIO3, chloric acid; also for hydrogen, — hydrocarbon, hydrogen and carbon 
 
 compound. 
 Hypo — ^ Prefix meaning under, less; hypochlorous acid less oxygen than 
 
 chlorous acid. 
 Hypothesis. A supposition; provisional view needing proof. 
 — ic. Termination meaning higher power of combination, more of the other 
 
 substance than ous. 
 
 — id (ide). Termination of name of- binary compound. 
 Inorganic. Mineral; not derived from animals or plants. See Organic. 
 Insoluble. Not soluble — in water, unless other solvent is mentioned. 
 
 — ite. See ate, 
 
 Kindling Temperature. Temperature necessary to start combustion. 
 
 Law. The statement, based on experiment, of any of Nature's regular actions 
 
 or conditions. 
 Iiitre. 1.06 quarts; 1000 cubic centimeters. 
 Metal. One of the heavier, lustrous, basic elements. 
 Metalloid. Non-metallic element. 
 Molecule. As commonly used, the relative weight of a compound represented 
 
CHEMICAL TERMS. lO/ 
 
 by formula; the smallest portion of a substance having the properties of the 
 
 substance ; a group of atoms chemically united. 
 Mon — Prefix, meaning one. 
 Monad. Element with valency one. 
 Monatomic. Supposed to have only one atom in its molecule; sometimes 
 
 used for univalent. 
 Mother liquor. The liquid from which a salt is, or has been, crystallizing. 
 Nascent. The condition of an element at the instant of release from chemical 
 
 union. 
 Neutral. Not affecting litmus ; neither acid nor alkaline. 
 Nomenclature. The system of naming (chemical substances). 
 Non-metal. One of the lighter acid-forming elements. 
 Normal Pressure, i atmosphere; 76om.m. 
 
 Normal Salt. A salt in which there is no (displaceable) hydrogen. NajS04. 
 Normal Solution. See page 69. 
 Normal Temperature. 0° Centigrade. 
 
 Normal Volume. Of a gas, when measured at, or calculated to, normal pres- 
 sure and temperature. 
 Notation. The system of viriting (chemical substances). 
 Organic. Product, direct or indirect, from animal or vegetable life; now- used 
 
 of carbon compounds. 
 — ous. Termination; see — ic. 
 
 Oxidize. Cause to combine with oxygen; also to raise to higher valence. 
 Ozy — . Containing oxygen; c g. oxysalt, oxyacid. 
 Per — Prefix, beyond; more; a greater quantity of — . 
 Precipitate. The finely divided solid formed by mixture of solutions. 
 Qualitative. Having reference only to the kind of matter. 
 Quantitative. Having reference only to the mass or weight or bulk of 
 
 matter. 
 Radical. A combination of elements which frequently occurs in compounds 
 
 NHj, SO4. 
 Reaction. A chemical change; action with indicators which show acidity, 
 
 alkalinity, neutrality; sometimes used for equation. 
 Reagent. Applied to the substance used to bring about a chemical change. 
 Reduce. To take away oxygen; to change to lower valence. 
 Saline. Consisting of, or belonging to, a salt. 
 Salt. The rather inactive compound resulting from a base and an acid, usually 
 
 solid and crystalline. 
 Saturated Compound. In which each element employs all its chemical 
 
 attraction; no place for more without either substitution or rearrangement. 
 
I08 EXERCISES IN CHEMISTRY. 
 
 Saturated Solution. Containing all the dissolved substance possible, at the 
 temperature. 
 
 Soluble. Capable of forming solution; unless otherwise mentioned, water is 
 the solvent. 
 
 Solution. A clear, uniform, liquid mixture; "homogeneous mixture which 
 cannot be separated into constituent parts by mechanical means." Ostwald. 
 
 Solution, Chemical. In which the action of the solvent makes a new com- 
 pound. 
 
 Solution, Physical. Applied to solution from which the original substance 
 can be recovered by evaporation. 
 
 Solvent. The liquid used in making a solution. 
 
 Speciiic Gravity. Times as heavy as an equal bulk of water. 
 
 Structural. Attempting to show the relation of the parts in a compound. 
 
 Sub — Prefix, under; less; lower. 
 
 Sublimation. Distillation applied to a solid. 
 
 Substitution. One substance displacing another from a compound. 
 
 Supersaturated. Containing more solid than is possible under ordinary con- 
 ditions. 
 
 Supporter of Combustion. Usually applied to oxygen. 
 
 Suspension. The turliid mixture of an insoluble solid with a liquid. 
 
 Symbol. The expression for an element and its relative weight. 
 
 Synthesis. Putting together to form more complex. 
 
 Ternary. A compound of three elements. 
 
 Theory. Supposition or explanation resting upon sound, but not complete, 
 evidence. 
 
 Tri — Prefix, three; threefold. 
 
 Uni — Prefix, like mono, one. 
 
 Valence. The number of times l of H or 35.4 of CI with which the atomic 
 weight of an element combines. See p. 65. 
 
 Vapor Density. The number of times as heavy as (the same bulk of) 
 hydrogen. 
 
 Volumetric. By volume or bulk measure. 
 
 Water of Crystallization. The definite proportion of water entering into 
 many crystals. 
 
 — yl. The termination of the name of nearly all radicals; (OH)' hydroxyl. 
 
 Practice in, and illustration of, the correct use of the fore- 
 going terms is a valuable exercise. For instance : — 
 
 Make correct statements illustrating the proper use of the 
 terms acid, alkali, amorphous, anhydrous, basicity, boil, 
 binary, etc. 
 
CHEMICAL TERMS. 
 
 Fill out the blank spaces correctly : — 
 
 Concentrated is 
 
 decomposes when heated. 
 
 is deliquescent. 
 Effervescence takes place when 
 
 is insoluble in water, etc. 
 
THE REFERENCE DIAGRAM. 
 
 The student of chemistry cannot hope to understand the 
 theoretical side of the subject until he has acquired a sound 
 basis of knowledge of its phenomena, laws, and terminology. 
 He must at first accept terms in their experimental meaning. 
 But later he must add prolonged and serious study of the 
 science as a whole or else remain forever in a mental fog 
 wherein premises and conclusions, phenomena and hy- 
 potheses, facts and their explanation are hopelessly mixed. 
 
 The accompanying Scheme, though not perfect, presents 
 to the eye the whole subject and the relations of its parts. 
 
 It is intended for a twofold use : — 
 
 (i) By advanced students as a basis for topical study ; 
 
 (2) By students in the elementary stage for frequent refer- 
 ence — to avoid misconceptions and to fix the relations of 
 one or another part to the whole. For instance, every 
 beginner must use the notation, but a glance at the diagram 
 shows that the notation of modern chemistry is the out- 
 growth of elaborate experiment and reasoning. 
 
 As the primary object is light upon the theories, small 
 space is given to applied chemistry. 
 
 The diagram represents present views and is therefore 
 only approximately historical. 
 
 In general the arrows start from data, sources, or premises 
 and point to results, inferences, conclusions. 
 
 The arrows (with a few exceptions) point in only one direc- 
 tion ; but every new conception, law, discovery, fulfilled pre- 
 diction which grows from these complex roots strengthens 
 every connected part. 
 
Diagram for Reference and Topical Study. 
 
INDEX. 
 
 Acetic acid, 20. 
 
 Acid, to litinus,.2o. 
 Advice to pupil, 17. 
 Air, 36. 
 
 oxygen in, 36, 82. 
 Alcohol, 31,46. 
 Alkaline, to litmus, 20. 
 Ammonia, 20, 30, 33, 37, 38, 43, 56. 
 
 dilute, 10. 
 Ammonium chlorid, 31, 36, 37. 
 compounds, 39. 
 nitrate, 39. 
 sulfate, 37. 
 sulfid, 12, 37. 
 Analysis of HCl, 88. 
 Antimony, 45, 50. 
 
 sulfid, 33, 45, 52. 
 Apparatus, names of, 16. 
 
 pupils', 9, 74, 94. 
 Arithmetic, 62-72. 
 Arsenic, 44. 
 Arsenious oxid, 27, 44. 
 Atomic Theory, diagram, 109. 
 in notation, 15. 
 place of, 6. 
 weight, 63. 
 weights, Table, 61. 
 Avogadro, Theory of, 62. 
 Barium chlorid, 21, 35, 41, 54, 78, 99. 
 peroxid, 25. 
 sulfate, 35, 41, 53. 
 Beef tea, 37. 
 Bleaching powder, 30. 
 Bone black, 47. 
 Borax, 45. 
 
 bead, 45. 
 Boric acid, 45. 
 Bromin, 31. 
 Bunsen burner, 50. 
 Cadmium, 51, ij. 
 
 chlorid, 39. 
 Calcium carbonate, 48. 
 chlorid, 27. 
 phosphid, 43. 
 sulfate, 28, 53. 
 Caoutchouc, 49. 
 Carbon, 47. 
 
 dioxid, 37, 48. 
 
 vol. of, 90. 
 
 Carbon, disulfid, 31, 32, 40, 42, 49. 
 monoxid, 48. 
 
 vol. of, 90. 
 Carbonates, 54. 
 Carbonic acid, 48, 
 Care of Chemicals, 8. 
 Charcoal, 33, 44, 47. 
 Chemical change by germs, 37. 
 equation, 14. 
 names, 15. 
 notation, 13. 
 properties, 20, 21. 
 union, 21. 
 Chemicals, care of, 8. 
 
 preparation of, 93, 
 Chlorids, 16, 52. 
 Chlorin, 29. 
 
 Classes of compounds, 16. 
 Combustion, 25, 50. 
 Compounds, names of, 16. 
 Computations, 60. 
 Copper, 20, 29, 34, 39, 44, 56, 87. 
 carboriate, 51. 
 nitrate, 40, 41, 56. 
 oxid, 51. 
 sulfate, 19, 21, 22, 27, 28, 51, 54, 
 
 79. 87- 
 Cupric chlorid, 29. 
 hydrid, 43. 
 Cuprous hydrid, 43. 
 Crystallization, 99. 
 
 water of, 27, 78, 100. 
 Secompositiou, 21. 
 
 double, 21. 
 Deliquescence, 27, loi. 
 Diagram, Reference, 109. 
 Distillation, 95. 
 Drying, 79, 97, 98, loi. 
 Efflorescence, 27, iol 
 Elements, list of, 61. 
 
 names of, 15. 
 symbol of, 13. 
 Equation, balancing, 65. 
 chemical, 14. 
 computations from, 68. 
 illustrated, 22. 
 reading, 17. 
 Equivalent of copper, 87. 
 lead, 87. 
 
112 
 
 EXERCISES IN CHEMISTRY. 
 
 Equivalent of magnesium, 76. 
 mercury, 87. 
 silver, 88. 
 tin, 77. 
 zinc, 76. 
 Erlenmeyer flask, 40, 97. 
 Etching on glass, 29. 
 Evaporation, 19, 100. 
 Examples, illustrative, 70. 
 Experiments, 19-56. 
 
 quantitative, 75-92- 
 Ferric chlorid, 52. 
 oxid, 20, 24. 
 Ferrous ammonium sulfate, ic2. 
 chlorid, 53. 
 oxalate, 51,90. 
 sulfid,3i, 33, 41. 
 Filtration, 19. 
 Flame, 50. 
 
 oxidizing, 51. 
 reducing, 51. 
 Formic acid, 48. 
 Formula, calculation of, 64. 
 
 computation from, 68. 
 meaning of, 13. 
 weight, 66. 
 GaseB, density of, 62. 
 
 formula of, 13, 64. 
 preparation of, 97, 
 purification of, 97. 
 volume of, 67. 
 weight of, 66. 
 Glacial phosphoric acid, 43. 
 Granulated zinc, 10. 
 Heat, source of, 9. 
 Hydrobromic acid, 31. 
 Hydrochloric acid, 22, 29, 30, 33, 44, 41;, 
 46, 48, 52, 76, 84, 88. 
 dilute, 10. 
 Hydrofluoric acid, 29. 
 Hydro-fluor-silicic acid, 47. 
 Hydrogen, 22. 
 
 dioxid, 28. 
 displaced, 75, -jf^. 
 phosphid, 43. 
 sulfid, 31, 33. 
 Hydrometer, 95. 
 Hydroxids, 52. 
 Hypophosphorous acid, 43. 
 Identification, by properties, 57. 
 Indicators, 83. 
 Ink, 30. 
 
 Insolubility, importance of, 55. 
 Insoluble compounds, 55, 98. 
 Introduction, 7. 
 lodin, 31. 
 Iron, 21, 23, 25. 
 
 Lead, 47, 87. 
 
 carbonate, 53. 
 
 nitrate, 33, 52, 53, 54, 56, 79, 91, 99. 
 
 oxid, 27, 47, 90. 
 
 sulfate, 29, 56, 91. 
 
 sulfid, II, 29, 33, 99. 
 Lime, 27, 37. 
 
 water, 27, 37, 48, 95. 
 Line of solubility, 81. 
 List for classification, 57-59. 
 
 of chemicals, 10, 75. 
 Logwood, 34, 47. 
 Lycopodium,5i.- 
 Magnesium, 51, 76, 88. 
 Manganese dioxid, 25, 28, 29. 
 Measuring apparatus, 80. 
 Mercuric chlorid, 56, 79, 87, 
 iodid, 102. 
 nitrate 51. 
 oxid, 21, 51. 
 Mercury, 47, 56. 
 Metals, precipitation of, 56, 87. 
 Metaphosphoric acid, 42, 43. 
 Methods of study, 8, 17, 57, 70, 74, 92, 
 
 93. 109- 
 Metric system and specific gravity, 69. 
 Molecular weight, calculation, 66. 
 
 determination, 62. 
 Names, Chemical, 15, 104. 
 Neutral, with litmus, 20. 
 Neutralization, 38. 
 
 quantitative, 83, 84. 
 Nitrates, 53. 
 Nitric acid, 20, 39, 41, 45, 86. 
 
 dilute, 10, 95. 
 Nitrogen, 36. 
 
 dioxid, 39. 
 
 monoxid, 39. 
 
 tetroxid, 40. 
 Normal solutions, 69, 94. 
 Notation, 13, 109. 
 Notes, 7. 
 Ozids, 51. 
 Oxygen, 24, 37, 40. 
 
 in air, 36, 82. 
 Ozone, 25. 
 Fara£B.n, 29. 
 Phosphoric acid, 43. 
 Phosphorus, 12, 36, 42, 82. 
 pentoxid, 42. 
 Physical properties, 20. 
 Potassium bromid, 31. 
 
 chlorate, 24, 76, 79. 
 
 chlorid, 24, 79. 
 
 ferricyanid, 56. 
 
 hydroxid, 84. 
 
 iodid, 26, 31, 54. 
 
INDEX. 
 
 "3 
 
 Potassium nitrate, 28, 41, 79. 
 nitrite, 36, 
 
 permanganate, 29, 34. 
 Preface, 5. 
 
 Preliminary manipulations, 19. 
 Preparation of gases, 97. 
 Preparations, 93. 
 Pressure, correction for, 67. 
 Processes, names of, 16. 
 Purification of gases, 97. 
 Pyrophosphoric acid, 43. 
 Qualitative Analysis, 33, 39, 55. 
 Quantitative experiments. 73-92. 
 Keaction, testing, 20, 83. 
 Reduction, 24, 43, 44. 
 Reference Diagram, 109. 
 Salts, foTmation of, 51-56. 
 preparation of, 99. 
 solution of, 27. 
 Silicon, compounds of, 46. 
 Silver, 33, 43, 87. 
 hydrid, 43. 
 
 nitrate, 25, 43, 53, 56, 87, 99. 
 Sodium carbonate, 26, 27, 28, 33, 39, 54, 
 79.83.85. 
 chlorid, 28, 30, 54, 79, 85, 88. 
 hydroxid, 37, 46, 51, 84. 
 hyposulfite, 28. 
 nitrate, 85. 
 phosphate, 35. 
 silicate, 46. 
 sulfate, 35, 54, 85. 
 sulfite, 34. 
 Solubility, importance of, 55. 
 line of, 81. 
 
 specimen table of, 55. 
 Soluble salts, 55. 
 Solution, 19. 
 
 approximate, 94. 
 exact, 69, 94. 
 of salts, 10, II, 12, 27 
 quantity in, 79. 
 
 Solution, saturated, 28, 95. 
 
 specific gravity of, 95. 
 supersaturated, 28. 
 Specific gravity, 69, 95. 
 
 heats, 64. 
 Starch, 11, 26, 32. 
 Sublimation, 44, 102. 
 Substitution, 22. 
 Sugar, 26, 35. 
 Sulfates, 53. 
 Sulfids, 52. 
 Sulfur, 20, 21, 26, 32. 
 Sulfuric acid, 22, 25, 30, 33, 34, 35, 38, 41, 
 
 47, 48, 88, 97, loi. 
 Supplies, 8, 74, 94. 
 Suspension, 19. 
 Symbol, 13. 
 
 Temperature, correction, 67. 
 Tension, correction, 67. 
 Terms, chemical, 104. 
 Theory, 109. 
 Tin, 41, 77. 
 
 chlorid, 39. 
 TTnion, chemical, 21. 
 Valence, 65. 
 Vapor density, calculation, 66. 
 
 determination, 62. 
 Volume experiments, 75, 77, 82, 88, 90. 
 Water, 26. 
 
 distilled, 26. 
 of crystallization, 27. 
 quantity, 78. 
 Weight of oxygen, 77. 
 
 lead nitrate, 91. 
 oxid, 90. 
 sulfate, 91. 
 dissolved salts, 79. 
 White lead, 90. 
 Wire gauze, 50, 
 Words, meaning of, 104. 
 Zinc, 10, 21, 22, 49, 76.