UC-NRLF 
 
 SB 35 M3E 
 
 LABORATORY DIRECTIONS 
 
 FOR 
 
 ELEMENTARv CHEMISTRY 
 
 MATT ILL 
 
jL^ J 
 
LABORATORY DIRECTIONS 
 
 FOR 
 
 ELEMENTARY CHEMISTRY 
 
 DESIGNED TO ACCOMPANY 
 
 A TEXTBOOK OF CHEMISTRY 
 BY W. A. NOYES 
 
 BY 
 HELEN ISHAM MATTILL, PH. D. 
 
 FORMERLY ASSOCIATE IN CHEMISTRY AT THE UNIVERSITY OP ILLINOIS 
 
 NEW YORK 
 
 HENRY HOLT AND COMPANY 
 1914 
 
COPYRIGHT, 1912, 1914 
 
 BY 
 HENRY HOLT AND COMPANY 
 
 IN MEMORIAM 
 
 THE. MAPLE- PRESS. YORK. PA 
 
PREFACE 
 
 The importance of the laboratory work in a course in elemen- 
 tary chemistry cannot be too strongly emphasized. It is only 
 by the actual contact and experimentation with some of the 
 many materials described in a text-book that the student ever 
 comes to any realizing sense of the means by which the science of 
 chemistry has been built up. For this reason it is advisable 
 to supplement the laboratory exercises with discussions of an 
 explanatory nature, in which the correlation between laboratory 
 and text-book work may be brought out. 
 
 The experiments described in this guide have been chosen 
 with the following objects in mind: To reduce the variety of 
 materials handled, consequently the confusion of many new 
 names and strange materials, to a minimum; to make each 
 experiment a fairly exhaustive study of some particular 
 material or property; to have the sequence of experiments 
 such as will lead to a certain amount of reasoning from analogy. 
 This guide does not pretend to be an exhaustive description of 
 elementary experiments, but rather a selection of a few suitable, 
 consecutive experiments from the many possible. 
 
 I wish to take this opportunity to acknowledge the many 
 helpful suggestions and criticisms which I have received from 
 Drs. C. W. Balke, C. H. Hecker, S. B. Hopkins and W. A. Noyes 
 in the preparation of this book. 
 
 in 
 
CONTENTS 
 
 PAGH 
 
 GENERAL DIRECTIONS 1 
 
 CHAPTER 
 
 I. APPARATUS AND MANIPULATION 10 
 
 II. PURE SUBSTANCES AND MIXTURES 16 
 
 III. OXYGEN 21 
 
 IV. HYDROGEN 26 
 
 V. WATER AND HYDROGEN PEROXIDE 31 
 
 VI. THE HALOGEN FAMILY 38 
 
 VII. SULFUR 55 
 
 VIII. NITROGEN 63 
 
 IX. THE ATMOSPHERE 68 
 
 X. PHOSPHORUS, ARSENIC, ANTIMONY AND BISMUTH 71 
 
 XI. CARBON 76 
 
 XII. SILICON 81 
 
 XIII. BORON 83 
 
 XIV. METALS AND NON-METALS 85 
 
 XV. THE ALKALI METALS 89 
 
 XVI. COPPER AND SILVER 92 
 
 XVII. CALCIUM, STRONTIUM AND BARIUM 95 
 
 XVIII. MAGNESIUM, ZINC AND MERCURY 98 
 
 XIX. ALUMINIUM 102 
 
 XX. LEAD AND TIN 104 
 
 XXI. CHROMIUM 107 
 
 XXII. MANGANESE 109 
 
 XXIII. IRON Ill 
 
 INDEX . 115 
 
LABORATORY DIRECTIONS 
 
 FOR ELEMENTARY 
 
 CHEMISTRY 
 
 GENERAL DIRECTIONS 
 
 Work in the laboratory should be undertaken with the fol- 
 lowing objects in mind: 1. By direct handling of the materials 
 to become familiar with certain "chemical reactrons. 2. By 
 analogy, and text-book study, to relate, the experiments 
 actually performed in the i3br3rai i or^ >> with jr 'as > 'maiiy facts 
 brought up in the lecture and quiz room discussion as possible. 
 3. By handling and manipulating apparatus to become familiar 
 with the tools at the disposal of the chemist, and the proper 
 use of the same, and at the same time to acquire ability and 
 dexterity in their use. A student, however good his under- 
 standing of chemical facts, has not reached the goal unless 
 he is able to perform neatly and properly any given chemical 
 manipulation. 4. To acquire habits of observation and a 
 proper discrimination between important and unimportant 
 details, and to draw conclusions of a general nature from a 
 specific case. 5. To record accurately, briefly, and promptly 
 the phenomena observed. 
 
 In the pursuit of these objects the keeping of a laboratory 
 note-book is of prime importance. It is absolutely essential 
 that these notes be recorded at the time the work is performed. 
 They should give a clear but brief account of work done and 
 should include a statement of the object of the experiment, a 
 record of all phenomena observed, such as change of color, 
 appearance of a precipitate or a gas, etc., and answer all ques- 
 
 1 
 
2 GENERAL DIRECTIONS 
 
 tions. A drawing of the apparatus used may often reduce 
 the length of the description required. Wherever possible a 
 chemical equation should be written to express the change 
 taking place. In order to write such equations it may be nec- 
 essary to refer to the text-book for the products of the reaction. 
 After the products have been determined the equation 
 should be written and balanced without reference to text- 
 book or other aid. Do not delay in recording the observations 
 made during a given experiment; each change should be jotted 
 down as observed. It is more essential that the note-book be 
 a record of the experiment taken at the time it is performed 
 than that it should be neat and well written, though both of 
 these ends may be accomplished at once without too great 
 difficulty. Under no circumstances should the record of an 
 experiment' jb, placed 'Qi) loose paper, the loss of which would 
 mean the repetition of the experiment so recorded. 
 
 The f pJlcwfiig, -gejner&l directions for manipulation should be 
 closely followed: 
 
 (1) Do not use too much material. The amount called for 
 in any experiment is ample to secure the results desired. Tak- 
 ing larger amounts of material is wasteful, frequently dangerous, 
 and always results in loss of time. 
 
 (2) Keep all bottles in their proper places. 
 
 (3) Material spilled on the side shelf should be promptly 
 swept into a waste jar. 
 
 (4) Never place the stopper of a reagent bottle on the table 
 hold it between the ringers. 
 
 (5) Do not dip stirring rods or tubes into the reagent bottles. 
 Pour a little of the reagent into a test tube. 
 
 (6) Do not waste gas or water. 
 
 (7) In pushing a glass tube through a stopper wet both the 
 tube and the stopper. It is wise also to protect the hand by 
 covering the tube with a towel. 
 
 (8) Do not heat beakers or flasks with the direct flame, but 
 always place them on a wire gauze. Test tubes are the only 
 glass receptacles which can be heated in the direct. flame. 
 
GENERAL DIRECTIONS 3 
 
 (9) Before leaving at the end of a laboratory period see that 
 all apparatus is in its proper place, and that the desk is clean. 
 
 DIRECTIONS FOR QUANTITATIVE EXPERIMENTS 
 
 In performing quantitative experiments the student must 
 keep in mind the object of the given experiment, and the fact 
 that successful results in quantitative work depend largely 
 on the care and accuracy with which given manipulations are 
 carried out. A given experiment may be successful in the 
 hands of one student and a complete failure in the hands of 
 another, not because of any fault in the choice of the experi- 
 ment, but because of the differences in manipulation which 
 result from an incomplete understanding of the object of the 
 experiment, and of the possibilities of error and how to avoid 
 them. For this reason it is especially essential that the direc- 
 tions for an experiment marked quantitative should be studied 
 from beginning to end before any manipulation is started. A 
 student should begin the manipulation of a quantitative 
 experiment with the object of the experiment so clearly in 
 mind that reference to the directions is unnecessary except as 
 regards the details of manipulation. 
 
 The record of a quantitative experiment should be kept in 
 tabulated form, and should in every case be kept in the note- 
 book. If the experiment has to be repeated the record of every 
 succeeding trial must also be in the note-book. On reading 
 over the directions make a note of what measurements are to 
 be taken and in what order they will be used, then prepare a 
 sheet of the laboratory note-book with the proper headings in 
 the proper order before starting the experiment. For example, 
 in the first experiment (No. 8) the weights of two tubes empty, 
 with potassium chlorate, and with the residue on heating potas- 
 sium chlorate are to be determined and from that data the 
 percentage of potassium chloride obtained on heating a given 
 amount of potassium chlorate is to be calculated. The headings 
 then might be arranged as follows: 
 
GENERAL DIRECTIONS 
 
 Small tube Large tube 
 
 Tube + potassium chlorate, 
 Tube empty, 
 Potassium chlorate, 
 
 Tube + potassium chloride (residue), 
 Tube empty, 
 Potassium chloride, 
 
 Per cent of potassium chloride in potassium 
 chlorate, 
 
 The arrangement of the data sheet before performing 
 the experiment is a good test of the student's understanding of 
 the object of the experiment, and for that reason should be 
 recorded in the note-book before the manipulation is begun. 
 
 While the directions for each quantitative experiment de- 
 scribe the special precautions to be observed in the given experi- 
 ment there are other precautions which apply to all quantitative 
 work and which should be kept constantly in mind, of which 
 the following are the most important. 
 
 Weighing. In all quantitative experiments use the chemical 
 balance. To prepare a dish for weighing it should be carefully 
 cleaned and wiped dry with a clean towel. After a dish has 
 been prepared for weighing it should be handled as little as 
 possible with the fingers because of deposition of moisture and 
 oil from the skin. Material may be placed in a dish so pre- 
 pared arid the required processes carried out according to the 
 directions of the given experiment, but the student should 
 always keep in mind the fact that an error may be introduced 
 into the weight of any given material by allowing foreign matter 
 to fall by accident into the dish, or by accidental loss of material 
 from the dish. Directions to avoid loss of material by spatter- 
 ing, etc., must be carefully observed. The success of every 
 quantitative experiment depends on determining by weighing, 
 measuring or otherwise, all of a given material which results 
 from a given chemical process. 
 
 Testing for Leaks. Some of the quantitative experiments 
 involve the evolution and measurement of a volume of gas. 
 
GENERAL DIRECTIONS 5 
 
 In all such experiments the apparatus must be constructed 
 with unusual care and tested to determine that there is no 
 possibility of a loss of gas through leaks. Rubber, instead of 
 cork stoppers should be used and wherever rubber tubing is 
 used in making connections the glass tubes, which should always 
 be fire polished (see Exp. 3 a), should meet inside the rubber 
 tubing. Unless an apparatus is air-tight it should not be used in a 
 quantitative experiment. 
 
 STANDARD CONDITIONS FOR GASES 
 
 In order that gas volumes may be comparable they must be 
 given under fixed conditions, since the volume of a given sample 
 of gas varies with both the temperature and the pressure under 
 which it exists. The fixed conditions which have been arbi- 
 trarily chosen by chemists are that the gas should be dry, 
 under the pressure of 760 mm. of mercury measured at C., 
 and at a temperature of C. The gas as obtained in the 
 laboratory is usually saturated with moisture, though in certain 
 cases, for example Exp. 20, it may be measured dry, and is 
 under the temperature and pressure conditions obtaining in the 
 laboratory. These conditions may be determined by a ther- 
 mometer and barometer. The barometer reading may indicate 
 a certain atmospheric pressure, but it is self-evident that if the 
 temperature of the room were lower the same atmospheric 
 pressure would give a lower barometric reading, because of the 
 contraction of the mercury on cooling. For this reason the 
 barometric reading must be corrected for the temperature at 
 which the barometer stands. A table of corrections for ordi- 
 nary temperatures can be readily prepared from the coefficient 
 of expansion of mercury, and may be found below. To correct 
 a barometric reading to C., then, the observed reading should 
 be reduced by the amount indicated in the table, for the 
 observed temperature. For instance suppose the barometer 
 reads 74.3 cm. at 17 C., the correction for 17 C. is 2.2 mm., 
 therefore the barometric reading at C. would be 743 2.2 = 
 
6 GENERAL DIRECTIONS 
 
 740.8 mm. But if the gas was saturated with moisture the 
 pressure indicated by the barometer was being divided between 
 the two components of the mixture, the water vapor and the 
 gas. Since the pressure of water vapor at varying temperatures 
 has been accurately determined, that may be obtained from a 
 table, and the pressure of the gas in the mixture is then the 
 difference between the barometric reading and the pressure 
 of water vapor at the observed temperature. For instance, the 
 corrected barometric reading, as found above, was 740.8 mm., 
 but the gas in question was saturated with moisture. The 
 vapor pressure of water at 17 C. is found from the table to be 
 14.5 mm., therefore the pressure of the dry gas was 740.8 14.5 
 = 726.3 mm. of mercury measured at C. The gas volume 
 must then be reduced to standard conditions of temperature 
 and pressure by applying the laws of Boyle and Charles, before 
 any comparison of volume may be made. 
 
 TABLE I 
 
 TEMPERATURE CORRECTIONS FOR REDUCING A COLUMN OF MERCURY 750 MM. 
 HIGH WHEN READ ON A GLASS SCALE, TO C., IN MILLIMETERS 
 
 Temperature, 15 16 17 18 19 20 21 22 23 24 25 
 Correction, 2 2.1 2.2 2.3 2.5 2.6 2.7 2.8 3.0 3.1 3.2 
 
 Temperature, 26 27 28 29 30 31 32 33 34 35 
 Correction, 3.4 3.5 3.6 3.7 3.9 4.0 4.1 4.3 4.4 4.5 
 
 TABLE II 
 VAPOR PRESSURE OF WATER IN MM. OF MERCURY 
 
 Temperature, 
 
 
 
 5 
 
 
 10 
 
 
 15 
 
 
 16 
 
 17 
 
 18 
 
 19 
 
 Pressure, 
 
 4.6 
 
 6, 
 
 5 
 
 9. 
 
 2 
 
 12. 
 
 8 
 
 13.6 
 
 14.5 
 
 15.5 
 
 16.5 
 
 Temperatures, 
 
 20 
 
 21 
 
 
 22 
 
 
 23 
 
 
 24 
 
 25 
 
 26 
 
 27 
 
 Pressure, 
 
 17.5 
 
 18. 
 
 .7 
 
 19 
 
 ,8 
 
 21. 
 
 1 
 
 22.4 
 
 23.8 
 
 25.2 
 
 26.7 
 
 Temperature, 
 
 28 
 
 29 
 
 
 30 
 
 
 31 
 
 
 32 
 
 33 
 
 34 
 
 35 
 
 Pressure, 
 
 28.4 
 
 30, 
 
 1 
 
 31, 
 
 ,8 
 
 33. 
 
 7 
 
 35.7 
 
 37.7 
 
 39.9 
 
 42.2 
 
 Temperature, 
 
 40 
 
 45 
 
 
 50 
 
 
 55 
 
 
 60 
 
 65 
 
 70 
 
 75 
 
 Pressure, 
 
 55.3 
 
 71. 
 
 9 
 
 92. 
 
 5 
 
 118. 
 
 1 
 
 149.5 
 
 187.6 
 
 233.8 
 
 289.3 
 
 Temperature, 
 
 80 
 
 85 
 
 
 90 
 
 
 95 
 
 
 100 
 
 
 
 
 Pressure, 
 
 355.5 
 
 433. 
 
 ,8 
 
 526 
 
 .0 
 
 634. 
 
 
 
 760 
 
 
 
 
GENERAL DIRECTIONS 
 
 TABLE III 
 
 DENSITIES AND PER CENT COMPOSITION OF WATER SOLUTIONS OF ACIDS 
 AND AMMONIA. (Taken from Landolt, Bornstein and Roth) 
 
 Acetic 
 
 acid 
 
 Nitric acid Phosphoric acid 
 
 Sp. Gr. 
 
 % acid 
 
 Sp. Gr. 
 
 % acid Sp. Gr. 
 
 %acid 
 
 = d 20 
 
 
 I 15 
 
 d 17 ' 5 
 
 
 4 
 
 
 4 
 
 4 
 
 
 1.012 
 
 10 
 
 1.06 
 
 10.67 1.06 
 
 10 
 
 1.026 
 
 20 
 
 1.12 
 
 20.22 1.12 
 
 20 
 
 1.038 
 
 30 
 
 1.19 
 
 30.87 1.18 
 
 30 
 
 1.049 
 
 40 
 
 1.25 
 
 39.8 
 
 .26 
 
 40 
 
 1.057 
 
 50 
 
 1.31 
 
 49.05 
 
 .34 
 
 50 
 
 1.064 
 
 60 
 
 1.37 
 
 59.36 
 
 .43 
 
 60 
 
 1.069 
 
 70 
 
 1.39 
 
 63.2 
 
 .53 
 
 70 
 
 1.070 
 
 80 
 
 1.42 
 
 69.77 
 
 .65 
 
 80 
 
 1.066 
 
 90 
 
 
 
 .76 
 
 90 
 
 1.050 
 
 100 
 
 
 
 
 Maximum density is at 
 78% acid. 
 
 Hydrochloric acid 
 
 Sulfuric acid 
 
 Sp. Gr. 
 
 %acid 
 
 Sp. Gr. 
 
 %acid 
 
 ^4 
 
 
 4 
 
 
 1.05 
 
 10.17 
 
 1.068 
 
 10 
 
 1.10 
 
 20.00 
 
 1.14 
 
 20 
 
 1.15 
 
 29.57 
 
 1.22 
 
 30 
 
 1.16 
 
 31.52 
 
 1.30 
 
 40 
 
 1.20 
 
 39.11 
 
 1.40 
 
 50 
 
 
 
 1.50 
 
 60 
 
 
 
 1.61 
 
 70 
 
 
 
 1.73 
 
 80 
 
 
 
 1.82 
 
 90 
 
 
 
 1.84 
 
 100 
 
 Ammonia 
 Sp. Gr. %NH 3 
 
 ,15 
 ' d "4 
 0.96 
 0.92 
 0.90 
 
 10 
 20 
 
 28.5 
 
8 GENERAL DIRECTIONS 
 
 TABLE IV 
 
 SOLUBILITIES. (Taken from Landolt, Bornstein and Roth.) Giving 
 in order the formula of the salt or its hydrate stable at 10, the weight in 
 grams of the salt, or the hydrate (always calculated from the formula given 
 in column 1) required to make a liter of normal solution, and the weight of 
 anhydrous substance in 100 grams of a saturated solution at 10, 20, 30, 
 40, 50 and 100. 
 
 Formula 
 
 Grams per 
 liter for a 
 normal 
 solution 1 
 
 Solubility 
 
 10 
 
 20 30 
 
 40 
 
 50 
 
 100 
 
 AgC 2 H 3 O2 
 AffTl 
 
 166.9 
 143.3 
 187.8 
 234.8 
 169.9 
 155.9 
 241.45 
 111.1 
 118.7 
 98.7 
 140.2 
 157.7 
 50. 
 109.5 
 37. 
 86. 
 114.1 
 154.3 
 119. 
 145.6 
 50. 
 147 9 
 
 0.875 
 8.9-10-5 
 
 1.037 
 1.53-10- 4 
 0.84-10- 5 
 3 53-10" 7 
 
 1.215 
 
 1.413 
 
 1.637 
 
 
 2.19-10-8 
 3.7-10- 4 
 
 AeBr 
 
 
 
 Agl 
 
 
 
 
 AgNOs 
 
 61.5 
 
 68.3 
 0.78 
 
 73. |77. 80. 
 
 30.1 
 1.46 
 
 Ag 2 SO4 
 A1C1 3 -6H 2 O ! 
 
 
 
 
 A1 2 (S04) 3 -18H 2 6.. 
 A1 2 K 2 (SO4)4-24H 2 O 
 BaCOs 
 
 25.1 
 
 26.6 
 
 28.8 
 
 7.74 
 
 31.4 
 
 34.3 
 
 47.1 
 30.6xaq. 
 
 1.8-10-3 
 25. 
 2.17 
 
 
 
 . 
 
 BaCl 2 -2H 2 O 
 
 26.3 ,27.6 
 3.36 4.75 
 
 29.- 
 6.85 
 
 30.4 '37. 
 10.5 
 2-10-s 
 
 BaO9H 2 O 
 CaCOs . . . . 
 
 CaCl 2 -6H 2 O 
 CaOH 2 O 
 
 39.4 
 
 42.7 
 0.123 
 0.20 
 
 53.54aq 
 0.113 0.104 
 0.21 0.21 
 56.3 57.51aq 
 58.4 ;61.4 
 36.1 39.4 
 
 61.42aq 
 0.096 0.05 
 
 CaS(V2H 2 6 
 CdCl 2 -2iH 2 O .... 
 Cd(NO 3 ) 2 -4H 2 O.... 
 CoCl 2 -6H 2 O 
 Co(NOs) 2 -6H 2 O.... 
 CrOs 
 Cu(NOs) 2 '6H 2 O 
 
 0.192 
 47.4 
 
 is'.s"" 
 
 59.5 
 o6'.7'2aq 
 
 
 31. 
 
 33.3 
 
 
 \fi9, h 
 
 64.6 
 
 67.4 
 
 
 55.6 
 47.9 
 
 51.6" 
 
 73.25aq 
 24.8 
 6.4 
 
 61.53aq 
 
 FeCl 3 -6H 2 O { 
 
 FeSO 4 '7H 2 O 
 H 3 B0 3 
 HgCl 
 
 90. i45. 
 
 75.9 
 
 28.7 
 8. 
 7-10~ 4 
 8.7 
 43 2 
 
 75.9 
 32.7 
 10.3 
 
 84. 2 an" 
 
 139. 
 21. 
 236. 
 135.6 
 119. 
 167. 
 87. 
 74.6 
 122.6 
 138.6 
 97 .2(64.8) 4 
 147. (49) 4 
 50.2 
 166. 
 214. 
 158. (31.6) 4 
 101. 
 87.2 
 101.7 
 128.3 
 123.3 
 99. 
 120.6 
 53.5 
 39.62 
 80.1 
 66.1 
 
 191.2 
 
 139. 
 151. 
 
 17. 
 
 21. 
 4.9 
 0.38 -10-* 
 6.9 
 39.4 
 6.54 
 
 28.2 
 
 
 HgCl 2 
 
 6.19 
 38. 
 
 
 
 10.2 
 
 35. 
 51.2 
 33.2 
 60.9 
 36. 
 35.9 
 15.8 
 44.2 
 50.5 
 
 KBr 
 
 KBrOs 
 
 
 11.7 
 
 53.9 
 
 28.7 
 12.7 
 
 54!s" 
 30. 
 16.5 
 5. 
 40.8 
 
 K 2 CO 3 -2H 2 O 
 
 53.3 
 
 27.2 
 9.2 
 
 KC1 
 
 23.8 
 
 4.7 
 
 25.5 
 
 6.8 
 1.9 (25) 
 
 KC1O 3 
 
 KC1O4 
 
 K 2 CrO 4 . . . 
 
 37.9 
 
 7.8 
 21.7 
 57.7 
 
 38.6 
 11.6 
 24.9 
 59.1 
 7.5 
 6. 
 24.1 
 10. 
 35.3 
 42. 
 26.2 
 
 39.5 
 15.4 
 28.1 
 60.4 
 10.5 
 8.3 
 31.6 
 11.5 
 
 40.1 
 22.6 
 31.2 
 61.5 
 11.4 
 10.4 
 39.2 
 13. 
 3fi 5 
 
 K 2 Cr 2 Or 
 
 KHCO 3 
 
 34.2 
 62.7 
 24.4 
 14.3 
 46.3 
 14. 
 
 KI... 
 
 67.6 
 
 KIOs 
 
 KMnO4 
 
 4.01 
 17.7 
 8.4 
 34.9 
 
 
 KNOs. . . 
 
 71.1 
 19.4 
 42.2 
 
 
 
 K 2 SO4 
 
 MgCl 2 -6H 2 O 
 Mg(NOs) 2 -6H 2 O. 
 MgS0 4 '7H 2 0.... 
 MnCl 2 -4H 2 O 
 MnSO 4 5H 2 O.... 
 NH 4 C1 
 
 45 . 9 
 
 
 23.6 
 
 29. 
 
 44 7 
 
 31. 
 
 33.56aq 
 49.5 
 37.3 
 33.5 
 
 42.5 
 53 . 7 2aq 
 24.9 
 43.6 
 
 
 38.6 
 27.1 
 
 17.4 
 
 43'." 
 
 39.4 laq 
 29.3 
 21.3 
 70.8 
 43.8 
 3 7 
 
 
 25. 
 13.7 
 
 42i2" 
 1.58 
 
 31.4 
 
 NH 4 HCO3 
 
 NH 4 N0 3 
 
 (NH4) 2 SO4 
 
 74.8 
 144.8 
 
 78.5 
 45.8 
 9 5 
 
 89.7 
 50.8 
 35. 5aq 
 55. an 
 47.6 
 
 Na 2 B407-10H 2 0. . 
 NaBr-2H 2 O 
 NaBrOs 
 
 47.5 
 
 27.7 
 
 
 51.4 
 33.4 
 
 53.7 
 
 
 
 
 
 
GENERAL DIRECTIONS 
 
 TABLE IV. Continued 
 
 Formula 
 
 Grams per 
 liter for a 
 normal 
 solution 
 
 Solubility 
 
 10 
 
 20 
 
 30 
 
 40 
 
 50 
 
 100 
 
 Na 2 CO3.10H 2 O... 
 NaCl 
 
 NaClOs 
 
 143. 
 58.5 
 106.5 
 171.2 
 149.1 
 42.2 
 85. 
 126. 
 161.2 
 124. 
 145.5 
 138.9 
 166. 
 73.8 
 133.3 
 141.8 
 132.8 
 91.8 
 91.1 
 148.8 
 143.8 
 
 11.2 
 26.3 
 
 17.6 
 26.4 
 49.7 
 44. 
 64. 
 8.8 
 46.8 
 20.5 
 16. 
 41. 
 49. 
 0.96 
 34. 
 1-10' 8 
 35. 
 41.5 
 0.7 
 15-10- 3 
 78.6 IJaq 
 54. 
 
 29. 
 26.5 
 
 47. 4aq 
 66.3 
 10. 
 49. 
 
 33.2 laq 
 26.7 
 56.5 
 49. 
 68.8 
 11.3 
 51. 
 
 32.2 
 26.8 
 
 31. 
 28.1 
 67.1 
 55 . 8 an 
 82. an 
 
 Na 2 CrOrlOH 2 0... 
 Na 2 Cr 2 07'2H 2 O . . 
 NaHCOs 
 
 33.4 
 63. 
 7.6 
 44.6 
 16.6 
 8.3 
 38. 
 
 51. 
 71.3 
 12.7 
 53. 
 
 NaNO 3 
 
 64.4 
 24 . 8 an 
 29.9 
 72.7 
 
 Na 2 SO 3 '7H 2 O 
 Na 2 S04'10H 2 O... 
 Na 2 S 2 O.r5H 2 O ... 
 Ni(\0 3 ) 2 -6H 2 O.. . 
 PbCl 2 
 
 29. 
 
 46. 
 
 32.5 an 
 50.6 
 55. 
 
 31.8 
 63. 2aq 
 
 
 
 3.2 
 56. 
 
 Pb(NO 3 ) 2 
 
 30.8 
 
 38. 
 
 41. 
 
 44. 
 
 SrCOs 
 SrCl 2 -6H 2 O 
 
 32.6 
 35.5 
 0.48 
 9.9-10-3 
 73.1 
 
 37 . 5 
 46.7 
 1. 
 
 40. 
 
 47.7 an 
 1.5 
 
 42.7 
 48.1 
 2.1 
 16- 10- 3 
 
 50.5 
 50.3 
 19.5 
 18-10-3 
 
 86. 
 
 44 ! 'laq ' 
 
 Sr(N0 3 ) 2 -4H 2 0.... 
 Sr(>9H 2 O 
 
 SrSO 4 
 
 ZnCl 2 -2 1/2 H 2 O . 
 Zn(NO 3 ) 2 -6H 2 O.... 
 ZnSO 4 -7H 2 O 
 
 
 82 an 
 
 
 67.4 3aq 
 41.2 
 
 43'.5'6aq 
 
 32.3 
 
 
 
 1 In many cases this quantity exceeds the solubility of the salt, but it is always given 
 because of the chemical significance of the value. 
 
 2 This is half the formula weight. 
 
 3 Anhydrous. 
 
 4 The figure in parenthesis applies when the salt is to be used as an oxidizing agent in 
 acid solution. 
 
 TABLE V 
 
 SOLUBILITY AND DENSITY OF GASES. (Taken from Landolt, Bornstein 
 and Roth.) Giving the volume of gas, measured at standard conditions, 
 absorbed by one volume of water, at 0, 20, 50 and 100 and the weight 
 per liter of the gas under standard conditions. 
 
 Gas 
 
 H 2 
 
 N 2 
 
 2 
 
 CL 
 
 HC1 
 
 Br 2 
 
 HBr 
 
 NH 3 
 
 NO 
 
 N 2 O 
 
 CO 2 
 
 S0 2 
 
 H 2 S 
 
 CH 4 
 
 C 2 H 4 
 
 C 2 H 2 
 
 
 
 20 
 
 50 
 
 100 
 
 0.0215 0.0184 
 0.0239 0.0164 
 0.0489 0.0310 
 4.61 2.26 
 506.9 442 3 
 
 0.0161 
 0.0106 
 0.0209 
 1.2 
 361 
 
 0.0160 
 0.0100 
 0.0170 
 0.000 
 
 60.5 21.3 
 612.5 
 
 6.5 
 
 468 
 
 
 
 1305 715.4 
 
 
 
 0.0738 0.047 
 0.629 
 
 0.0315 
 
 0.0262 
 
 1.713 0.878 
 79 8 39 4 
 
 0.436 
 
 
 
 4.6 2.5 
 0.0556 0.033 
 0.226 0.122 
 
 1.4 
 0.0213 
 
 0.8 
 0.017 
 
 1 . 73 1 . 03 
 
 
 
 Weight 
 per liter 
 0.0898 
 1.256 
 1.429 
 3.180 
 1.610 
 7.143 
 3.616 
 0.762 
 1.340 
 1.970 
 1.977 
 2.870 
 1.542 
 0.717 
 1.252 
 1.162 
 
CHAPTER I 
 APPARATUS AND MANIPULATION 
 
 Check up the apparatus found in the desk, using a list 
 obtained from the storeroom. Record any missing articles on 
 a sheet furnished for that purpose, and after getting the approval 
 of an instructor take both lists to the storeroom and obtain the 
 missing articles. 
 
 Study the apparatus supplied with a view to determining its 
 uses and proper manipulation as follows: 
 
 1. Bunsen Burner. 
 
 a. Examine the construction of the Bunsen burner. Explain 
 the function of each part. 
 
 b. Attach the burner to the gas supply by means of a rubber 
 tube, open the holes at the bottom of the vertical tube and light 
 the gas. What is the character of the flame? Show the 
 structure of the flame by a sketch. Explore the various parts 
 of the flame with a platinum wire. What are the relative 
 temperatures of the different parts? Where should an object 
 be placed to secure the maximum heating effect? Could a 
 flame of this character be used for illuminating purposes? 
 What is the principle of the Welsbach light? 
 
 c. Close the holes at the bottom of the tube. What is the 
 character of the flame now? Hold a cold porcelain dish in the 
 flame for a moment. What is the deposit? (Illuminating gas 
 contains a compound of carbon and hydrogen.) Why is it 
 possible to obtain a deposit of this substance on a cold dish? 
 Hold a clean portion of the dish above, but not in, the flame. 
 Why is no deposit obtained in this case ? What causes the 
 luminosity of the flame? 
 
 10 
 
BUNSEN BURNER 
 
 11 
 
 d. Repeat c using a non-luminous flame. Explain the 
 result. 
 
 e. Extinguish the flame, turn on the gas without lighting it, 
 and hold the dish in the current of gas. Why is no deposit 
 obtained? 
 
 /. Light the gas and show the structure of the luminous flame 
 by a sketch (turn the flame low to avoid flickering). Test the 
 temperature of this flame with a platinum wire. How does it 
 compare with the temperature of the non-luminous flame ? 
 
 g. Quickly introduce the head of a match into the innermost 
 cone of both flames. What does this experiment indicate as to 
 
 Burrette 
 
 7 
 
 Pipette 
 FIG. 1. 
 
 FIG. 1. 
 
 the temperature of the central cone? What composes this 
 central portion? Devise and perform an experiment to test 
 this point. 
 
 2. Measuring Instruments. 
 
 The measuring instruments consist of the cylinder, pipette 
 and burette (side shelf). Fig. 1. The cylinder may be filled to 
 the zero mark, then the amount removed read on the scale. The 
 
12 APPARATUS AND MANIPULATION 
 
 pipette is used to measure one stated quantity. When the 
 drawn out tip is placed under the surface of the liquid and 
 suction is applied to the other end until the bottom of the men- 
 iscus of the liquid is brought to the level of the mark on the 
 neck, the pipette will deliver the stated quantity. 
 
 The burette is filled to the zero mark, taking care that no 
 bubbles of air remain in the parts about the nozzle, the desired 
 quantity of liquid withdrawn from the nozzle at the bottom, 
 and the volume withdrawn read on the scale. All readings 
 should be taken from the bottom of the meniscus. 
 
 Fill a test-tube with 20 c.c. of water from the pipette, mark 
 the level of the water on the tube with a paper label or file 
 mark. Determine the capacity of the tube up to the given 
 mark by means of the cylinder and the burette. Which gives 
 the most accurate measure of volume, and why? 
 
 3. Working with Glass Tubing. 
 
 a. To cut glass tubing make a file mark at the desired spot, 
 then grasping the tube with both hands place the two thumb 
 nails directly behind the mark and break by pressing out and 
 pulling at the same time. The sharp edges of the freshly 
 broken glass must be rounded off by holding it in the hottest 
 part (Exp. 1) of the Bunsen flame till it begins to glow, then 
 allowing it to cool slowly. This is called fire polishing. 
 
 6. To make a bend in glass tubing it is desirable to heat the 
 tubing uniformly through a distance of 5 to 8 cm. (2 to 3 
 inches), so as to avoid too great thickening of the glass at the 
 bend. This is accomplished by holding the glass parallel with 
 the flame of a fish tail burner, and just above the inner cone, 
 i.e., in the hottest part of the flame. The tubing should be 
 slowly rotated in the flame until it has softened, then withdrawn 
 from the flame, bent to the desired angle, and allowed to cool 
 slowly. Fig. 2. 
 
 c. To draw out a fine capillary, or a jet, the glass should be 
 softened for only 1 to 2 cm. of its length. What flame is best 
 
GLASS TUBING 13 
 
 adapted to this purpose? Rotate the glass tube in the flame 
 until softened uniformly, then remove it from the flame and 
 pull it out to the desired length. 
 
 d. After a little practice in making bends and jets make a 
 wash bottle as shown in Fig. 3. A two-hole rubber stopper or 
 a cork stopper may be used. If a cork is used it should be 
 softened by means of a press or cork roller and the holes must 
 
 i ii 
 
 FIG. 2. FIG. 3. 
 
 I. Shows bends made by heating the glass over a very short distance, as 
 would be done by a Bunsen burner. The glass buckles and is uneven in 
 thickness. II. Shows similar bends made by heating the glass over a 
 greater distance, using a fish-tail burner. The glass is even in thickness. 
 
 be made by cutting through with a cork borer which is slightly 
 smaller in diameter than the glass tubing. Give the cork and 
 borer opposite rotatory motion, with a steady pressure not great 
 enough to tear the cork, until it is cut half way through, then 
 remove and start from the other end, placing the borer so that 
 the cuts will meet in the middle. Always remove the cork from 
 inside the borer. If the hole in the cork is too small for the 
 glass tubing it may be enlarged by using a rat-tail file. 
 
14 APPARATUS AND MANIPULATION 
 
 4. The Balance. 
 
 a. Under the direction of an instructor examine the chemical 
 balance. Sit directly in front of the center of the balance and 
 release the stop with a slow, steady movement. If the pointer 
 does not swing raise the stop and release again. The swing of 
 the pointer should be entirely within the graduated portion of 
 the ivory scale. When the pointer is swinging freely, find the 
 center of the arc described in its swing. This is the true zero 
 and its position must be determined for each weighing, since the 
 point changes. 
 
 b. Place on the left hand pan of the balance (see rules in 
 paragraph c) one of the hard glass tubes which has been previ- 
 ously cleaned and dried (see page 4). With the aid of the 
 forceps place on the right hand pan that combination of weights 
 which is nearest to, but less than, the weight of the tube. Al- 
 ways begin with the largest weight and work down. When the 
 tube is nearly counterpoised close the balance case and adjust 
 the rider upon the scale beam until the pointer oscillates about 
 the true zero, as determined in a. Find the weight of the tube 
 first by adding together the values of the weights which have 
 been removed from the set; and second by adding the weights 
 themselves. If these two weights are identical, record in the 
 note-book. Notice the decimal system in the weights. For 
 example if a dish is balanced when the opposite pan contains 
 10 grams, 2 grams, 500 mg., 200 mg., 100 mg., 10 mg., and the 
 rider is upon 2 on the beam, the weight of the dish is 12.812 
 grams. Now remove the weights, replacing each in its proper 
 position in the set, raise the rider from the beam, and close 
 the balance case. 
 
 c. Always observe the following rules in weighing: 
 
 (1) Handle the balance with care. 
 
 (2) Raise the stop gently, and always as the pointer is passing 
 the zero mark. 
 
 (3) Never add anything to the pans, or remove anything 
 from them, while they are swinging. 
 
 (4) Never place chemicals directly upon the pan. Use a 
 
THE BALANCE 15 
 
 piece of glazed paper or a watch glass which has been properly 
 counterpoised. 
 
 (5) Always use forceps in handling the weights. 
 
 (6) Place the object to be weighed and the weights as near 
 the center of the pan as possible. 
 
 (7) Do not allow the pans to swing. 
 
 (8) In case any solid is spilled in the balance case use the 
 camels hair brush to remove it. Keep the balance and the shelf 
 clean. 
 
 (9) Keep both eyes open while reading the swing of the 
 pointer. 
 
 (10) If the balance is not in proper adjustment, call an 
 instructor. 
 
CHAPTER II 
 
 PURE SUBSTANCES AND MIXTURES 
 
 5. Boiling Point as a Test for Purity. 
 
 This test depends on the fact that a pure substance has a 
 fixed boiling-point, while a mixture of two or more substances 
 has a boiling-point which changes as the boiling proceeds. In 
 such a case the substance which has the lower boiling-point 
 boils off in greater proportion at first, and if the vapors are con- 
 
 FIG. 4. 
 
 densed a liquid can be obtained which contains a higher per- 
 centage of the low boiling substance than the original mixture. 
 This fact is utilized in many commercial processes, as in the 
 manufacture of alcohol. 
 
 a. Connect a 50 c.c. distilling flask with a Liebig condenser. 
 Fig. 4. Force a slow stream of cold water through the jacket 
 
 16 
 
BOILING POINT 17 
 
 of the condenser in the direction of the arrows. Why is the 
 water circulated in this direction? Half fill the flask with dis- 
 tilled water and stopper with a one-hole cork carrying a 
 thermometer. The bulb of the thermometer should be just a 
 little below the level of the delivery tube of the flask. Why? 
 Heat the distilling flask with a small flame so that the water 
 distils slowly. Collect the distillate in a small flask. After a 
 few cubic centimeters of water have been collected read the 
 temperature. Does this stay constant as the distillation pro- 
 gresses? Is it the temperature given in the text-book as the 
 boiling-point of water? If not, give the reason for the differ- 
 ence. What ought to be specified when a boiling-point is given? 
 
 b. Add 5 c.c. of alcohol to 20 c.c. of water and repeat Exp. a 
 with this mixture. What is its boiling-point? Do alcohol and 
 water form a mixture or a pure substance? Collect the first 
 few cubic centimeters of the distillate in an evaporating dish and 
 apply a flame to the liquid. In a like manner collect a few 
 cubic centimeters after about half of the original mixture has 
 been distilled, and als.o the last few cubic centimeters. Do not 
 heat the distilling flask to dryness. Test each portion of the 
 distillate w r ith the flame. Account for the difference. Is there 
 any alcohol in the last portion of the distillate? Is there any 
 water in the first portion? 
 
 c. Dissolve a few small crystals of copper sulfate (CuS0 4 ) in 
 25 c.c. of water and distil the mixture until about 5 c.c. have been 
 collected. What is the distillate? From what kind of im- 
 purities can water be freed by distillation? Evaporate a little 
 of the distillate on a watch glass, holding the watch glass about 
 6 inches above the tip of the bunsen flame. Evaporate tap 
 water similarly. Why is distilled water used in the laboratory? 
 
 6. Melting-point as a Test for Purity. 
 
 This test depends on the fact that a pure substance has a 
 fixed melting-point, while a mixture of two substances begins 
 to melt at a temperature different from that at which it is 
 completely melted, or, in reverse order, begins to solidify at a 
 
18 PURE SUBSTANCES AND MIXTURES 
 
 temperature different than that at which it is completely 
 solidified. A familiar example is found in the melting of ice. 
 Pure ice melts at C., and ice and water can exist together 
 only at C. If salt (an impurity) is added to water, the water 
 begins to freeze at some temperature below C., and the mass 
 is not completely solid (frozen) until 23 C. is reached. To 
 illustrate this test proceed as follows: 
 
 a. Place about 2 grams of powdered, hydrated calcium chloride 
 (CaCl 2 '6H 2 0) in a clean, dry test-tube and heat very gently 
 (this salt has a very low melting-point) until the substance 
 softens enough so that the bulb of a thermometer may be 
 inserted into it. Continue warming, while stirring with the 
 thermometer, until a clear liquid is obtained. Now let the 
 melted mass cool gradually, while continuing the stirring, by 
 holding the test-tube under running water, and watch the 
 thermometer constantly. When the solid particles first begin 
 to appear note the temperature, and continue the cooling till 
 the contents of the tube is solid. Has the temperature changed 
 during the process of solidification? Has the substance a fixed 
 freezing-point? Is it a pure substance? If the first attempt 
 is not successful this procedure may be repeated as many 
 times as necessary. 
 
 b. While the thermometer is still in the tube warm again, this 
 time noting the temperature at which the substance melts. 
 (Caution: If heat is applied too rapidly the melting-point 
 cannot be observed; the heating should be so gradual that the 
 solid is completely melted only after several minutes.) Is 
 the melting-point the same as the freezing-point? 
 
 c. Now add to the tube and contents a small amount of solid 
 sodium chloride (NaCl) and proceed as above to determine the 
 freezing-point. Has the substance a fixed freezing-point or 
 melting-point now? Account for the difference. 
 
 7. Elements and Compounds. 
 
 a. Grind 4 grams of sulfur and 7 grams of iron filings to- 
 gether in a mortar. Examine the mass with a lens. Place a 
 
ELEMENTS AND COMPOUNDS 19 
 
 small amount of the mixture in each of three test-tubes. To the 
 first add 15 c.c. of water and shake thoroughly. To the second 
 add 10 c.c. of dilute hydrochloric acid. To the third add 5 c.c. 
 of carbon disulfide, shake, filter 1 (use a dry filter) the liquid 
 into a watch crystal and allow it to evaporate spontaneously. 
 (Carbon disulfide is inflammable. Do not bring it near a 
 flame.) Describe what takes place in each case. 
 
 b. Heat the remainder of the original mixture in a test-tube 
 until it begins to glow strongly. Allow the mass to cool, reduce 
 it to a powder, and repeat the tests made in a. Account for 
 the differences in the results obtained. 
 
 c. Heat a small portion of mercuric oxide in a dry, hard glass 
 test-tube. Watch for the appearance of the evolution of a gas, 
 and test it with a glowing splinter. Note the change in the 
 appearance of the residue in the test-tube, and the deposit on 
 the cool parts of the glass. Is mercuric oxide an element or 
 compound? Why? Have you any experimental proof for 
 saying whether the products of this decomposition are elements 
 or compounds? 
 
 d. Heat a small portion of copper oxalate in the same way. 
 Test the gas evolved with a glowing splinter, and note the change 
 in appearance of a rod wet with lime water and held at the 
 mouth of the tube. Is the same gas evolved as before? What 
 remains in the tube? Continue the heating until the residue 
 in the tube shows a further change. Is this change accompanied 
 by the evolution of a gas? 
 
 e. Define an element, and a compound. Show which materials 
 used above have been proved to be compounds. If the weight 
 of the residue from the copper oxalate had been determined 
 before and after the final change from the dark red to the black 
 substance, the change could have been shown to be accom- 
 
 1 To prepare a filter paper fold it through the middle, then fold the straight 
 edge onto itself, pinching the cone so made together just at the fold. Now 
 open the cone so that there are three thicknesses of paper on one side and one 
 on the other; fit into the funnel, taking care that the paper lies against the 
 glass wall on all sides, finish creasing the second fold in the paper, and 
 moisten so that it will stick in place. 
 
20 PURE SUBSTANCES AND MIXTURES 
 
 panied by increase in weight. Is it then a decomposition or a 
 combination? 
 
 8. Law of Definite Proportion. Quantitative. 
 
 Read the Directions for Quantitative Experiments (page 3) 
 before proceeding with this experiment. 
 
 a. Weigh accurately the two hard glass ignition tubes or 
 use those weighed in Exp. 4. In the smaller place about 0.5 
 gram, in the larger about 1 gram of pure, dry potassium 
 chlorate (KC10 3 ) (balance room) and weigh again. The differ- 
 ence between the two weights is the exact weight of the 
 potassium chlorate. Heat each tube gently in the Bunsen 
 flame. The potassium chlorate will melt, and then appear to 
 boil. During the boiling the heating should be just sufficient 
 to sustain boiling, as too great heat will cause the loss of solid 
 particles along with the escaping gas. A glowing splinter should 
 be applied to the gas escaping from the tube, the result to be 
 explained later. After the boiling has ceased the tubes should 
 be heated with the hottest part of the Bunsen flame till the 
 material is completely melted, then allow them to cool and w^eigh 
 again. Repeat the heating and weighing until constant weight 
 is obtained. The difference between the last weight and that 
 of the empty tube gives the weight of the solid residue resulting 
 from the decomposition of potassium chlorate by heat. Keep 
 these tubes with tfie residues until your results have been approved 
 by an instructor. 
 
 6. Calculate the per cent of residue in each case and compare. 
 State the law of definite proportion and show how this experi- 
 ment supports the law. 
 
 c. This example is a case of the decomposition of a more com- 
 plex compound (potassium chlorate) into an element (oxygen) 
 and a simpler compound (potassium chloride). Write the equa- 
 tion. How was the oxygen recognized? Dissolve the residue, 
 after the results of the quantitative work have been approved, and 
 test the solution with a solution of silver nitrate (AgN0 3 ). 
 Apply the same test to a solution of some of the original po- 
 tassium chlorate. Interpret the results. 
 
CHAPTER III 
 
 OXYGEN 
 
 9. Catalysis. 
 
 a. Melt a small sample of potassium chlorate (KC10 3 ) in a 
 dry test-tube and note the rate at which decomposition proceeds; 
 then drop into the melted mass a pinch of manganese dioxide 
 (Mn0 2 ) (known to be free from organic matter) and note the 
 change in the rate of decomposition. Allow the mass to cool 
 and dissolve. Has there been any apparent change in the 
 manganese dioxide? 
 
 6. Place as much sodium peroxide (Na 2 2 ) as can be piled on 
 a dime in each of two porcelain crucibles; to one of the crucibles 
 add a very small amount of copper oxide (CuO), then fuse the 
 material in each crucible, allow to cool and to each add water and 
 note the rate of evolution of gas. Test the gas evolved with a 
 glowing splinter. What is it? Does the copper oxide have any 
 effect on the nature of the chemical change? On the speed? 
 Has it undergone any change during the reaction? Boil the 
 mixture in the crucible containing the copper, and filter. The 
 copper oxide (CuO) was converted into the hydroxide (Cu(OH) 2 
 or CuO-H 2 0) by the fusion, and on boiling changed back to 
 copper oxide. 
 
 10. Preparation and Properties of Oxygen. 
 
 a. Calculate from the data of Exp. 8, and from the ca- 
 pacity of a gas-collecting bottle and the weight of 1 liter 
 of oxygen, the amount of potassium chlorate (KC10 3 ) re- 
 quired to give oxygen enough to fill four of the gas-collecting 
 bottles. Weigh out the calculated amount of potassium 
 chlorate on the rough laboratory balance, mix with one-fourth 
 its volume of manganese dioxide (Mn0 2 ) and place in the large 
 
 21 
 
22 OXYGEN 
 
 hard glass test-tube which should be fitted with a stopper 
 carrying an outlet tube, and which must be proved to be air-tight 
 before proceeding. A small portion of this mixture should be 
 tested by heating in another tube. If the reaction takes place 
 quietly it is safe to conclude that organic matter is absent and 
 the reaction will proceed quietly in the large tube. Connect 
 the delivery tube with a glass tube leading to the trough in 
 which the gas-collecting bottles have been filled with water 
 and inverted (Fig. 5) and gently heat the potassium chlorate 
 till gas is evolved, collecting the gas in the inverted bottles. 
 After the bottles have been filled with gas a glass plate should 
 
 FIG. 5. 
 
 be placed over the mouth while still under water, then the 
 bottle and cover may be removed and set aside for future use. 
 Always withdraw the delivery tube from below the surface of 
 the water before allowing the generating tube to cool. Why? 
 The supply of the gas may be stopped almost instantaneously 
 at any time by withdrawing the heat from the generating tube. 
 Test the four bottles of the gas as follows, covering the bottles 
 again as soon as the spoon is withdrawn. 
 
 b. Into one insert an iron wire or picture cord, the tip of 
 which has been heated and dipped into sulphur, then lit. 
 
 c. Into the next insert some burning sulphur on the iron 
 spoon. 
 
DENSITY OF OXYGEN 23 
 
 d. Into the third insert a spoon carrying some burning 
 red phosphorus. 
 
 e. Into the fourth insert some burning magnesium ribbon, 
 held with the pincers. 
 
 /. Add about 10 c.c. of water to each of the last three bottles 
 and shake, then test the solutions with both blue and red 
 litmus paper, testing at the same time a sample of the water 
 added. What has been the nature of the reaction in these 
 cases? Why did combustion cease before all the material had 
 burned? Are the products of combustion soluble in water? 
 To what can you ascribe the difference in the effect on litmus 
 paper? Classify these three elements according to the effect 
 of the solution of their oxides on litmus. Are they at the same 
 time classified with respect to any physical properties? Why was 
 it necessary to test a small portion of the mixture of potassium 
 chlorate and manganese dioxide by heating separately? What 
 would be the result if combustible matter, as charcoal, had been 
 present in the mixture? 
 
 11. The Effect of Concentration on the Speed of a Reaction. 
 
 Will iron wire burn in the air as it does in oxygen? Burn 
 sulfur, phosphorus and magnesium in bottles filled with air. 
 Note the rate of combustion as compared with burning in 
 oxygen and account for the difference. Add water and test 
 with litmus as in Exp. 10 /. Are the products of the reaction 
 the same in the two cases? 
 
 12. To Determine the Weight of a Liter of Oxygen. Quan- 
 titative. 
 
 Various methods for determining the density, or weight per 
 liter of a gas, may be used. In the following experiment the 
 weight of the gas will be determined by the loss in weight of a 
 material, due to evolution of the gas in question, while the 
 volume of the gas will be measured directly. Having the 
 weight of a given volume the weight per liter can then be calcu- 
 lated by direct proportion. 
 
24 
 
 OXYGEN 
 
 a. Fit the large ignition tube with a rubber stopper and de- 
 livery tube bent so that it can be connected with a gas burette 
 (Fig. 6). Fill the burette with water and see that there is no air 
 in the connecting rubber tubing by raising and lowering the 
 reservoir (R) several times, while the pinch cock is open. Con- 
 nect the burette with the ignition tube and test for air tightness. 
 Then remove the ignition tube, weigh it, add 0.2 to 0.5 gram of 
 potassium perchlorate (KC10 4 ) and weigh again. Bring the 
 level of the water in the burette near the zero mark, attach the 
 
 FIG. 6. 
 
 ignition tube and open the cock. If the level of the water 
 in the burette remains constant for two minutes (showing the 
 apparatus to be air tight) bring the gas to atmospheric pressure 
 by raising or lowering the reservoir until the level of the water 
 in the two tubes is the same, then record the burette reading. 
 Now heat the potassium perchlorate gently until about 30 c.c. 
 of gas has been evolved, then allow to stand till the tube has 
 reached room temperature. Read the volume of the gas in 
 the burette after bringing the level of the water in the reservoir 
 even with that in the burette (the cock should remain open 
 
DENSITY OF OXYGEN 25 
 
 through all this procedure), read the temperature and baro- 
 metric pressure in the laboratory, then remove and weigh the 
 ignition tube and residue. If moisture has collected in the 
 ignition tube allow it to stand over night, or until it has com- 
 pletely dried by spontaneous evaporation, before weighing. 
 The difference between the two burette readings is the volume 
 of oxygen, the difference between the weights of the tube and 
 potassium perchlorate and the tube and residue is the weight of 
 the oxygen. 
 
 b. Since the volume of a gas varies greatly under varying con- 
 ditions of temperature and pressure, it is convenient to state the 
 weight of a given volume under " standard" conditions, i.e., 
 at C. and 760 mm. pressure. Calculate the volume the 
 oxygen would occupy under standard conditions, remembering 
 that the gas measured in the burette is a mixture of oxygen 
 and water vapor, and that the atmospheric pressure read on 
 the barometer must be corrected for temperature (see page 5). 
 The pressure correction to be applied for both water vapor and 
 the temperature of the barometer may be found in tabulated 
 form on page 6. From the weight of the given volume of 
 oxygen calculate the weight of one liter. Calculate the per 
 cent by which your result differs from the accepted value, 
 1.429 grams. Keep the ignition tube and residue for further 
 work. (Exp. 24.) 
 
 c. Why is it necessary to allow for the vapor pressure in the 
 gas? Why does not the fact that the ignition tube was full 
 of air, some of which has gone with the oxygen into the gas 
 burette, vitiate the results? How would it affect the result 
 if the ignition tube had not cooled completely before reading 
 the volume of oxygen? Why must the water in the reservoir 
 be brought to the level of that in the burette before reading 
 the volume of oxygen? 
 
CHAPTER IV 
 
 HYDROGEN 
 13. Preparation and Properties of Hydrogen. 
 
 a. Construct an apparatus consisting of a 250 c.c. generating 
 flask (the Erlenmeyer is well adapted to this purpose) stoppered 
 with a cork carrying a thistle tube reaching to within 0.5 cm. 
 of the bottom of the flask and a delivery tube bent for the 
 
 FIG. 7. 
 
 delivery of gas into bottles inverted in the pneumatic trough 
 (Fig. 7). This apparatus must be air tight and may best be 
 tested for tightness by closing the open end of the delivery 
 tube with a rubber tube and pinch cock, then pouring water 
 into the thistle tube. Explain why the water cannot run down 
 into the flask if the apparatus is tight. 
 
 If the following directions are observed carefully there will 
 
 26 
 
PREPARATION OF HYDROGEN 
 
 27 
 
 be no danger of an explosion, though as further protection a 
 towel may be wrapped about the generating flask, if desired. 
 6. Place about 5 grams of granulated zinc in the generating 
 flask, pour in enough dilute hydrochloric acid to cover the 
 lower end of the thistle tube and collect, the gas by holding an 
 inverted test-tube over the end of the delivery tube. Carry the 
 test-tube mouth downward to the Bunsen burner, which must 
 be at least two feet from the delivery tube, repeating the process 
 until the gas in the test-tube burns quietly. Try the effect if 
 
 FIG. 8. 
 
 the test-tube of gas is carried mouth upward to the flame. 
 What information does this give regarding the density of the gas? 
 Collect a bottle of hydrogen by displacement of water and 
 while holding it mouth downward insert into it a burning 
 splinter. Why does the flame go out, and why is the splinter 
 ignited again as it is withdrawn from the tube ? 
 
 c. Attach a glass jet to the delivery tube by means of a short 
 piece of rubber tubing. Collect a test-tube of hydrogen from 
 this jet by displacement of air (should the test-tube be inverted 
 or upright?) and after lighting the test-tube of gas at the Bunsen 
 
28 HYDROGEN 
 
 flame carry it back quickly to the jet. If the hydrogen is 
 pure enough to be lighted with safety it will light in this way. 
 Never apply a match to a hydrogen generator, but always light it 
 with a test-tube of the burning gas as described. Hold a watch 
 glass above the flame for a moment. What is the deposit? 
 Explain its formation. 
 
 d. Remove the glass jet and replace it with a hard glass tube 
 open at each end and containing about 0.5 gram copper oxide 
 (CuO), Fig. 8. After again testing the apparatus to make sure 
 it is free from an explosive mixture warm the copper oxide 
 gently while passing hydrogen over it. Hold a watch glass 
 at the outer end of the combustion tube. What is deposited? 
 What remains in the combustion tube? What has been 
 reduced during this action? What has been oxidized? Write 
 the reaction. 
 
 e. Clean out the hard glass tube and repeat d using magnetic 
 oxide of iron in place of cupric oxide, and observing all pre- 
 cautions to avoid an explosion. Write the equation for this 
 reaction. What has been oxidized during this action? 
 
 /. May other metals and acids be used in the preparation of 
 hydrogen? Test this by pouring dilute sulfuric acid on sam- 
 ples of magnesium, aluminium and copper, each in a separate 
 test-tube. The gas evolved may be tested by bringing the 
 mouth of the test-tube to the flame. Add a few drops of 
 dilute nitric acid to a sample of lead in a test-tube. Does the 
 gas evolved have the properties of hydrogen? In all these 
 cases if the reaction is not spontaneous the test-tube and its 
 contents should be warmed gently before drawing any 
 conclusions. 
 
 14. The Rate of Diffusion of Gases. 
 
 Draw out the closed end of a test-tube to a capillary (if the 
 closed end gets too hot to hold, a piece of glass tubing may be 
 sealed to it by heating the two together for a moment) cut the 
 capillary, and reduce the size of the opening by heating, gently, 
 until when the tube full of air is inserted open end down into a 
 
DIFFUSION OF GASES 
 
 29 
 
 beaker full of water (Fig. 9) it takes about one-half minute for the 
 water to rise in the tube to a mark made about two inches 
 from the open end. The water in the beaker must of course 
 come above this mark on the test-tube. In order that the fol- 
 lowing measurements may be comparable the test-tube should 
 always be brought to the bottom of the beaker, tipped just 
 enough so there is free access of water from the outside, and 
 the level of the water in the beaker should be always the 
 same. Now determine accurately the time elapsing between 
 the moment when the tube is inserted into the beaker of water 
 and when the water comes to the mark on the test- tube, 
 repeat this several times with the tube full of air, recording 
 each determination, then fill the tube with 
 hydrogen by displacement of air and again 
 determine several times the time required 
 for the water to rise to the mark, as before. 
 Remember that in order to completely fill 
 this tube with the easily diffusible hydro- 
 gen a rapid stream of the gas should be 
 led into it, and in order to keep the hydro- 
 gen after the tube is filled the thumb may 
 be placed over the open end, but the time 
 elapsing between filling the tube and insert- 
 ing it into the water must be as brief as pos- 
 sible. Repeat, filling the tube with carbon dioxide from a 
 Kipp generator, observing all the precautions as in the case of 
 hydrogen. Calculate the densities of these gases from their 
 rates of diffusion, taking hydrogen as 1, and compare with the 
 values given in the table. 
 
 15. The Combining Weight of Copper. Quantitative. 
 
 a. Fit up the hydrogen generator as described in Exp. 13, and 
 connect the delivery tube with a hard glass combustion tube 
 open at each end, held horizontally by the iron clamp, and from 
 which a right angle bend leads to the bottom of a test-tube, 
 Fig. 8. After proving the apparatus to be air tight disconnect 
 
 FIG. 9 
 
30 HYDROGEN 
 
 the combustion tube, place in it about 0.5 gram (do not weigh 
 accurately) of powdered copper oxide (CuO), and again connect 
 with the generator. Place about 7 grams of granulated zinc in 
 the generator and add dilute hydrochloric acid till the lower 
 end of the thistle tube is covered. See that the delivery tube is 
 at least two feet away from any flame. Collect a test-tube of the 
 gas by displacement of air, and carry it, inverted, to the flame, 
 repeating the process until the gas in the test-tube burns 
 quietly. Why? Now heat the copper oxide gently, having the 
 delivery tube bent down and leading to the bottom of a test- 
 tube. Notice the change in the copper oxide, and continue 
 the heating until no further change takes place (this will require 
 about 20 minutes), and all the water formed has been driven 
 out of the combustion tube into the test-tube, then allow the 
 tube and contents to cool while hydrogen is still passing through. 
 (Be sure that the evolution of hydrogen continues through all 
 the heating.) Write the equation for the reaction. What is 
 the reducing agent in this action? What is the oxidizing agent? 
 
 6. If the copper oxide has been completely reduced, and the 
 product resulting allowed to cool in an atmosphere of hydrogen, 
 it should be pure copper. A weighed amount of this copper 
 may now be converted to the oxide, weighed, and from these 
 two weights the weight of oxygen in the oxide is obtained by 
 difference. The combining weight of copper, or the weight of 
 copper which will combine with 8 grams of oxygen, may be 
 calculated by proportion. 
 
 In order to do this, clean and weigh a porcelain crucible. 
 Place in it part of the copper obtained above, and weigh again. 
 Now heat the crucible and contents until the red copper color 
 has entirely disappeared, stirring occasionally with the platinum 
 wire. Allow the crucible to cool, and weigh it. Heat it again 
 and weigh again, until two successive weighings are within 
 2 mg. the same. Calculate the combining weight of copper. 
 
 16. The Equivalent Weight of Magnesium, Exp. 93, page 98, 
 may be Performed here if Desired. 
 
CHAPTER V 
 
 WATER AND HYDROGEN PEROXIDE 
 17. Composition of Water. 
 
 a. The qualitative composition of water by synthesis has 
 been illustrated. Review the experiments. 
 
 6. The qualitative composition of water by analysis may be 
 illustrated by the decomposition of water vapor by hot iron filings. 
 
 Fit a 100 c.c. flask with a stopper carrying a thistle tube and 
 an obtuse angle delivery tube, which should be connected with a 
 
 FIG. 10. 
 
 hard glass combustion tube, the further end of the combustion 
 tube being connected with a delivery tube leading under 
 water, Fig. 10. Make all connecting tubes as short as possible 
 so as to avoid condensation of the steam. After making sure 
 that the apparatus is air tight place about 50 c.c. of water in 
 the flask and about a gram of iron filings in the combustion tube. 
 Clamp the tube so that it slopes down toward the flask slightly 
 (to prevent water from condensing and running over the hot 
 
 31 
 
32 WATER 
 
 part of the glass) and heat the iron filings with one Bunsen 
 burner, at the same time bringing the water in the flask just to 
 boiling with another burner. For successful results the iron 
 must be kept at a low red heat while steam is passing slowly 
 over it. Collect the decomposition product of the steam by 
 displacement of water, and test for hydrogen. What was 
 the source of the hydrogen? Where will you expect to find 
 the oxygen? Remove some of the material from the com- 
 bustion tube. Does it appear the same as the iron filings used? 
 Refer to the text-book for the composition of the iron oxide, 
 and write the equation. 
 
 Compare this equation with the one written in Exp. 13 e. 
 What is the relationship between these two equations? What 
 kind of a reaction is this? What causes the equilibrium to 
 shift in opposite directions in these two experiments? 
 
 18. Solubility of Salts in Water. Quantitative. 
 
 The solubility of a salt is defined as the weight of salt in 
 100 grams of a saturated solution at the given tempera- 
 ture. The solubility determination is usually made by weigh- 
 ing a given amount of solution known to be saturated at a 
 given temperature, removing the water by spontaneous 
 evaporation and finally by heating, then weighing the 
 dry salt. 
 
 a. Cover each of two evaporating dishes with a watch glass or 
 glass plate and weigh the dish with cover. 
 
 b. Prepare a saturated solution of potassium dichromate 
 (K 2 Cr 2 7 ) or potassium chloride (KC1) by adding the powdered 
 salt to about 50 c.c. of warm (40) water until some remains 
 undissolved after stirring intermittently for ten minutes. The 
 success of this experiment depends on the preparation of a truly 
 saturated solution, therefore plenty of time must be allowed for 
 all of the salt that will to dissolve; the temperature must be kept 
 as constant as possible, the solution must be stirred, and the 
 salt must be powdered so that it will dissolve rapidly. Pour 
 about 10 c.c. (estimate the quantity, do not measure it) of this 
 
SUPERSATURATED SOLUTIONS 33 
 
 saturated solution into one of the weighed evaporating dishes, 
 noting the temperature of the solution at the moment of pouring. 
 Weigh the solution in the covered dish immediately, and only 
 to the second decimal place, i.e., record the weight as soon as 
 you come to the place where 0.01 gram more is too much. Then 
 place the dish in a safe place in the desk and remove the cover, 
 leaving it to evaporate spontaneoulsy, or hasten the evaporation 
 by placing the dish on the steam bath. In the meantime 
 the saturated solution will have cooled down, and, after 
 stirring vigorously and allowing the crystals to settle, a 
 second sample should be placed in the second weighed dish, 
 the temperature at the time of pouring noted, and the pro- 
 ceeding repeated. 
 
 c. After the solutions in the dishes appear to have evaporated 
 to dryness replace the covers and heat each one gently with the 
 Bunsen burner, keeping the cover on to avoid loss from decrepi- 
 tation of the crystals. Decomposition will result if potassium' 
 dichromate is heated above its melting-point. If moisture 
 collects on the cover warm until it has entirely disappeared, 
 then allow to cool and weigh. 
 
 d. From the data calculate the weight of potassium dichro- 
 mate or potassium chloride which will dissolve in 100 grams of 
 water at the two temperatures. Look up the solubility of the 
 salt used at 10 and 20, and draw a solubility curve from these 
 four points. 
 
 19. Supersaturated Solutions. 
 
 a. Define a solvent, a solute, saturated, unsaturated and 
 supersaturated solutions. Could a solution in contact with 
 the solid solute remain permanently unsaturated or supersatu- 
 rated? How would you make a supersaturated solution of a 
 salt more soluble in hot water than in cold? Of a salt more 
 soluble in cold water than in hot ? Give two examples of unsta- 
 ble equilibrium and show in what respect each differs from 
 stable equilibrium. 
 
 b. To 25 c.c. of water add 20 grams of hydrated sodium sul- 
 fate (Na 2 S0 4 '10H 2 0) and warm with stirring until solution is 
 
34 WATER 
 
 complete. Decant into a perfectly clean beaker, cover with a 
 watch glass and set aside, allowing the solution to cool spon- 
 taneously. Agitating the solution, or the presence of any solid 
 matter, may cause crystallization to take place during the cool- 
 ing. When the solution has cooled to room temperature, if no 
 crystallization has taken place, drop in one crystal of sodium 
 sulfate, and explain the phenomenon. If crystallization 
 occurs during cooling the crystals will have to be redissolved 
 by warming again, and the cooling and subsequent treatment 
 with a crystal of sodium sulfate repeated. What is the condi- 
 tion of the solution (saturated, unsaturated or supersaturated) 
 (1) when poured into the clean beaker, (2) when the undisturbed 
 solution has cooled, (3) after crystallization has taken place? 
 What test could be applied to determine whether a given solu- 
 tion is saturated, supersaturated or unsaturated? 
 
 20. Water of Hydration. 
 
 Heat 1 gram samples of copper sulfate, potassium nitrate, 
 sodium chloride, sodium sulfate and potassium dichromate 
 each in a separate test-tube, and determine which of them con- 
 tain water of hydration. Place the samples of copper sulfate 
 and sodium sulfate, after they have been heated in the test- 
 tube, on watch glasses, and on two other watch glasses place 
 samples of the same materials which have not been heated. 
 Label these four samples and allow them to stand over night 
 in the desk, then examine. In which case has the change 
 brought about by heating been reversed? To what is this 
 reversal due? Which of the salts is efflorescent and which 
 deliquescent? Describe a means by which a deliquescent salt 
 may be made to effloresce. 
 
 21. Natural Waters. 
 
 a. Evaporate about 5 c.c. each of distilled and tap water on 
 watch glasses on the water bath, note the quantity of the residue 
 in each case and test it to see if it is again soluble in water. 
 
 b. Test portions of tap and distilled water for chlorides by 
 adding a few drops of dilute C.P. nitric acid (HN0 3 ) (side shelf) 
 
HYDROGEN PEROXIDE 35 
 
 and silver nitrate (AgN0 3 ) solution. (A white precipitate in- 
 dicates the presence of chlorides.) Test other portions for 
 carbonates by adding freshly filtered lime water (a white precipi- 
 tate indicates the presence of carbonates) and for calcium salts 
 by adding to fresh portions ammonium hydroxide (NH 4 OH) 
 and ammonium oxalate ((NH 4 )2C 2 04) solution (a white preci- 
 pitate shows the presence of calcium salts) . Compare the tests. 
 c. If rain water and sea water are available repeat these tests 
 on them, and compare all four waters. 
 
 HYDROGEN PEROXIDE 
 
 22. Preparation and Properties of Hydrogen Peroxide. 
 
 a. To make a solution of hydrogen peroxide add gradually 
 about 1 gram of sodium peroxide (Na 2 2 ) (carry this in a 
 watch glass or evaporating dish, as it will cause spontaneous 
 combustion of organic material such as paper) to about 50 c.c. 
 of water in a flask, which should be kept cool by holding under 
 running water. Why? Then add dilute sulfuric acid, gradu- 
 ally and with stirring, until litmus paper placed in the solution 
 just turns red, then test the solution as follows (save some of this 
 solution for Exp. 23) : 
 
 6. Prepare some lead sulfide by passing hydrogen sulfide 
 through 1 c.c. of a solution of lead acetate (Pb(C 2 H 3 2 ) 2 ), or 
 by adding a few drops of acetic acid to 1 c.c. of lead acetate, 
 then adding 0.5 c.c. of ammonium sulfide. Add a few cubic 
 centimeters of the hydrogen peroxide solution to this mixture, 
 warm, and note the change in color of the precipitate. What 
 is the reaction? 
 
 c. In another test-tube place 1 c.c. of a solution of potassium 
 permanganate (KMnOJ, add a few drops of dilute sulfuric 
 acid, and then the hydrogen peroxide. Test the escaping gas 
 with a glowing splinter. What is it? Write the equation. 
 
 23. To Determine the Concentration of the Hydrogen Peroxide. 
 Quantitative. 
 
 Fill a gas burette with water and attach a short-stemmed 
 funnel to the rubber tubing above the stopcock (Fig. 11). Bring 
 
36 
 
 WATER 
 
 the water up to the top of the graduated tube and close the 
 cock, then introduce into the funnel 10 c.c. of the hydrogen 
 peroxide solution prepared in Exp. 22, measured by the pipette. 
 By opening the cock carefully the solution can be drawn into 
 the gas burette without admitting any air after it. If air is 
 admitted it must be forced out by raising the reservoir and 
 opening the cock, but the mixing of the solution with the water 
 in the burette which this causes is not desirable. Now place 
 about 20 c.c. of potassium permanganate solution, which has 
 been acidified with 2 c.c. of dilute sulfuric acid, in the funnel 
 and draw about 2 c.c. at a time into the burette, until the color 
 of the potassium permanganate in the burette is permanent, 
 
 remembering that the solution can be 
 drawn in only when the level of the 
 water in the outer reservoir is below 
 that in the burette, and do not at any 
 time draw in all the solution from the 
 funnel. When the reaction has ceased, 
 as indicated by the permanence of the 
 color, read the volume of oxygen lib- 
 erated at atmospheric pressure, the 
 barometer and the thermometer. Cal- 
 culate the weight of oxygen obtained 
 (remember that it was saturated with 
 moisture). Write the equation for the 
 reaction involved. What per cent of 
 the oxygen liberated comes from the 
 hydrogen peroxide? From what source 
 does the remaining oxygen come? 
 Calculate the weight of the hydrogen peroxide which was present 
 in the 10 c.c. portion of the solution. Assuming the solution 
 to have a specific gravity of 1, calculate the per cent by weight 
 of hydrogen peroxide in the solution. 
 24. Law of Multiple Proportion. Quantitative. 
 
 In experiment 8 (the law of definite proportion) the amount of 
 potassium chloride in potassium chlorate was determined. 
 
 FIG. 11. 
 
LAW OF MULTIPLE PROPORTION 37 
 
 The amount of potassium chloride in potassium perchlorate may 
 be determined in a similar fashion, and from these two deter- 
 minations the amount of oxygen united with a given amount of 
 potassium chloride in each of the compounds calculated. If 
 the ignition tube and residue from Exp. 12 have been kept, the 
 residue (the partially decomposed potassium perchlorate) 
 should be heated until there is no further evolution of oxygen, 
 then cooled and weighed. From the weights taken in Exp. 12. 
 and this final weighing, calculate the per cent of potassium 
 chloride in potassium perchlorate. 
 
 From the data obtained in Exp. 8 calculate the amount of 
 oxygen which is combined with 1 gram KC1 in KC10 3 . From 
 the data in this experiment calculate the amount of oxygen 
 which is combined with 1 gram KC1 in KC10 4 . What is the 
 ratio of these two amounts? 
 
 State the law of multiple proportion and show how this 
 ratio illustrates it. 
 
CHAPTER VI 
 
 THE HALOGEN FAMILY 
 
 25. Preparation of Chlorine. 
 
 Caution: Do not allow chlorine to escape into the air un- 
 necessarily, as the fumes are poisonous and cause very unpleasant 
 colds if breathed in small quantities. 
 
 By what general method may chlorine be prepared from hy- 
 drochloric acid ? What sort of a reaction would this be ? What 
 general class of agents would be used for this purpose? Wliat 
 would you expect to be the result if materials which give 
 oxygen on heating were warmed with hydrochloric acid? 
 Test your conclusion by placing very small amounts of 
 potassium chlorate (KC10 3 ), potassium perchlorate (KC10 4 ), 
 and manganese dioxide (Mn0 2 ) in separate test-tubes and 
 warming with about 1 c.c. of dilute hydrochloric acid. The 
 gas evolved should be tested for chlorine by holding a strip 
 of paper wet with a starch potassium iodide solution (made 
 by mixing equal volumes of starch solution and potassium 
 iodide (KI) solution) above the mouth of the tube. If chlor- 
 ine is present the paper will turn blue, due to the formation 
 of a compound of starch and iodine which is blue, but a large 
 amount of chlorine may further change the color to brown, 
 then bleach it. The brown color is that of free iodine, and 
 the bleaching is due to the oxidation of iodine by chlorine 
 and water, forming iodic and hydrochloric acids. Write the 
 reactions. As soon as the test has been made pour the con- 
 tents of the tube down the sink, and follow with plenty of 
 water. 
 
 26. The Density of Chlorine. Quantitative. Poison. 
 
 The method used in determining the weight of a liter of 
 oxygen cannot be used in this case because of the solubility 
 
 38 
 
DENSITY OF CHLORINE 
 
 39 
 
 of chlorine in water, and the lack of a material which can be 
 used to give off chlorine as potassium perchlorate gives oxygen. 
 The method used for chlorine in this experiment consists in 
 weighing the same volume of air and of chlorine under known 
 conditions, and finding the volume weighed by replacing the 
 chlorine with a liquid, whose volume can be measured. The 
 volume of the chlorine is then equal to the volume of the 
 liquid necessary to replace it, and its weight is equal to the 
 
 FIG. 12. 
 
 weight of that same volume of air + the difference between 
 the weight of the flask and air, and the flask and chlorine. 
 
 a. Carefully clean and dry a 250 c.c. receiving flask, fit it with a 
 cork stopper and weigh the flask and stopper, recording the atmos- 
 pheric pressure and temperature at the same time. Now fit 
 up a chlorine generator similar to that used for hydrogen, 
 taking special care to have all rubber connections as short as 
 possible, since chlorine attacks the rubber. The delivery tube 
 from the generator should lead, through a cork, to the bottom 
 
40 THE HALOGEN FAMILY 
 
 of a Woulff bottle, the other neck of which is fitted with a 
 cork carrying a delivery tube which projects just below the 
 surface of the cork, and is bent so that it will reach to the 
 bottom of the receiving flask (Fig. 12). Place about 4 grams 
 of manganese dioxide (Mn0 2 ) in the generating flask and add 
 10 c.c. of concentrated hydrochloric acid, diluted with 10 c.c. 
 of water. The chlorine coming from this generator should be 
 dried by bubbling through concentrated sulphuric acid in the 
 Woulff bottle (a depth of 1 cm. is sufficient) and should then 
 be led to the bottom of the receiving flask to prevent diffusion 
 into the air as far as possible. Generate the chlorine by 
 warming gently and as soon as a green gas appears in the 
 generating flask introduce the delivery tube into the weighed 
 flask, leaving the flask unstoppered while chlorine passes in. 
 As soon as the flask seems to be full of chlorine, as indicated 
 by the color of the gas, remove the delivery tube and stopper 
 the flask. Collect the rest of the chlorine in three gas-collecting 
 bottles the mouths of which are closed as nearly as possible 
 with glass plates, cover, and set aside for Exp. 27. 
 
 b. Weigh the stoppered flask containing the chlorine, again 
 noting the atmospheric pressure and temperature, remove the 
 cork and invert the flask over a beaker containing 225 c.c. 
 of water and 25 c.c. of potassium hydroxide (KOH) solution. 
 Do not allow the flask filled with chlorine to stand any length 
 of time before this absorption, as the chlorine is absorbed by 
 the cork, ruining the experiment. The chlorine will dissolve 
 in the alkaline solution, causing the liquid to be sucked up into 
 the flask. When the reaction seems to be complete place the 
 beaker and flask, with its mouth still under the surface of the 
 solution in the beaker, in a sink full of water, lower the beaker, 
 bring the level of the liquid inside the flask even with the water 
 outside, and while holding the flask in this position put in the 
 stopper. Now remove the flask, and measure the volume of 
 the solution it contains, which is equal to the volume of air 
 replaced by chlorine. 
 
 c. The weight of the air replaced by chlorine may be calculated 
 
CHLORINE 41 
 
 from its volume, the temperature and pressure at the first 
 reading (assume that the air was dry) and the weight of a 
 liter of air under standard conditions. The difference between 
 the weight of that volume of air and of chlorine may be found 
 by subtracting the weight of the flask and air from the weight 
 of the flask and chlorine, and that difference, plus the weight 
 of the air ; gives the weight of chlorine occupying the given 
 volume under the second set of conditions. From this calculate 
 the weight of a liter of chlorine. 
 
 27. Preparation and Properties of Chlorine. Poison, 
 
 a. If a supply of chlorine has not already been prepared in the 
 previous experiment (26a.) it may be prepared by placing about 
 4 grams of manganese dioxide (Mn0 2 ) in a generator similiar 
 to that used for hydrogen, adding through the thistle tube 10 
 c.c. of concentrated hydrochloric acid diluted with 10 c.c. of 
 water, and warming gently. The delivery tube should reach to 
 the bottom of a gas collecting bottle (should the bottle be up- 
 right or inverted?) which should be covered as completely as 
 possible with a glass plate and the bottle should be replaced as 
 soon as the greenish color of the chlorine becomes marked. 
 Fill three bottles in this manner, cover and set aside and im- 
 mediately place the delivery tube of the generator under a 
 dilute potassium or sodium hydroxide solution; force the 
 chlorine out of the generator and through the alkaline solution 
 (what is the reaction between chlorine and the alkali?) by pour- 
 ing water through the thistle tube until the generator is filled, 
 then empty out the generator before proceeding. 
 
 b. Scatter a little powdered antimony into one of the bottles 
 of chlorine Write the reaction. 
 
 c. Place strips of moist and of dry colored calico in one bottle 
 and account for the difference in behavior. 
 
 d. Drop moist pieces of red and blue litmus paper into 
 the other, then introduce a jet of burning hydrogen (from a 
 Kipp generator) until the appearance of the flame indicates 
 that the chlorine is all used up, then again drop in pieces of 
 
42 THE HALOGEN FAMILY 
 
 moist litmus. To what property of chlorine is the effect on 
 litmus due? Why must the litmus be moist? What is the 
 product when hydrogen burns in chlorine, and what is its 
 effect on litmus? Would you expect to get the second test 
 with litmus if any free chlorine was left in the gaseous mixture? 
 Give the reason for your answer. 
 
 28. Hydrochloric Acid. 
 
 a. Review the experiment in which hydrochloric acid was 
 prepared by synthesis. 
 
 b. Calculate the amounts of sodium chloride and sulfuric acid 
 required by the equation NaCl + H 2 S0 4 = NaHS0 4 + HCl to 
 give enough hydrochloric acid to fill two gas-collecting bottles, 
 and to saturate 5 c.c. of water at 20 C. Place the calculated 
 amount of sodium chloride in a generating flask fitted with 
 a thistle tube and L-shaped delivery tube and add through 
 the thistle tube a little more than the calculated amount of 
 concentrated sulfuric acid (sp. gr. 1.84), diluted with one-half 
 
 its volume of water. In handling concentrated 
 sulfuric acid be careful not to let it touch the 
 skin, as it causes unpleasant burns; if any 
 does get on the skin it should be washed off 
 at once with plenty of water. In diluting sul- 
 furic acid always pour the acid into the 
 water. Never pour water into concentrated sul- 
 furic acid. 
 
 Place the generator on the wire gauze and 
 warm gently, and with the delivery tube ex- 
 tending to the bottom of the gas collecting 
 bottle, which should be dry, collect two bottles of the gas 
 by displacement of air. The bottle is full of hydrochloric 
 acid when moist blue litmus paper held just above the neck 
 of the bottle turns red. Cover these bottles and set aside 
 for future use. Then connect the delivery tube with a U 
 tube (storeroom) as shown in Fig. 13. Place 5 c.c. of water 
 in the U tube and continue the generation of hydrochloric acid 
 
HYDROCHLORIC ACID 43 
 
 till the water is saturated. Save the generator and contents 
 for future work. 
 
 c. Blow across the mouth of one of the bottles. Explain the 
 formation of the fumes. Introduce into the same bottle a 
 glass rod wet with ammonium hydroxide. What are the 
 fumes in this case? 
 
 d. Invert the other bottle over water in the sink. What prop- 
 erty of hydrochloric acid does this illustrate? If any gas is 
 left in the bottle after shaking about with the water, test it 
 to determine whether it is hydrochloric acid or air. Explain 
 the test and results. 
 
 Why not collect the hydrochloric acid by displacement of 
 water? How would you determine whether a solution of 
 hydrochloric acid is saturated or not? What method should 
 be used to determine the density of hydrochloric acid and 
 why? 
 
 e. Place 1 c.c. of the solution from the U tube in a test-tube, 
 and add a granule of zinc. What is the reaction? Pour the 
 solution from the excess of the zinc and evaporate to dryness 
 in an evaporating dish. What is the residue? 
 
 /. Add portions of the solution to a small amount of calcium 
 oxide, until the latter is entirely dissolved. Evaporate this 
 solution to dryness in an evaporating dish, and allow it to 
 stand over night in the dish and examine. What is the 
 reaction? What causes the change on standing? 
 
 g. Add a few drops of silver nitrate (AgN0 3 ) solution to 
 another sample of the hydrochloric acid. What is formed? 
 Where was this test used before? For what radical is it a 
 test? 
 
 h. What are the products when hydrochloric acid acts on 
 (a) metals, (b) oxides, (c) hydroxides, (d) oxidizing agents 
 (Exp. 25)? 
 
 i. Pour the contents of the generating flask, after cooling, into 
 a beaker containing about 10 c.c. of water (or if the contents 
 of the flask solidify on cooling add 10 c.c. of water to the 
 flask) and after allowing time enough for the solution to become 
 
44 THE HALOGEN FAMILY 
 
 saturated (fifteen minutes at least) pour some of the solution 
 from the solid residue and add a few cubic centimeters of 
 concentrated hydrochloric acid. How does this experiment 
 prove the reaction reversible? What means may be used 
 to force the reversible reaction to completion? 
 
 29. lonization. 
 
 a. Test the conductivity of each of the following by placing 
 about 10 c.c. in a U tube, and inserting the platinum wires 
 from the battery into each arm. If a current flows ; the am- 
 meter which is connected "in series" with the battery will 
 indicate it by the position of the pointer. Rinse the platinum 
 wires in distilled water between each test. 
 
 (a) Distilled water. 
 
 (b) Dilute hydrochloric acid. 
 
 (c) Dilute potassium hydroxide (KOH) solution. 
 
 (d) Dilute sodium chloride (NaCl) solution. 
 
 (e) Alcohol and water. 
 
 (f) Dilute sugar solution. 
 
 Notice the phenomena which take place at the poles when a 
 current flows. Test to determine which pole gives hydrogen and 
 which oxygen or chlorine. Which pole is the anode? What 
 happens to the sodium ion when it loses its electrical charge 
 at the cathode, when electrolyzing sodium chloride solution? 
 Under what condition could metallic sodium be prepared by 
 electrolysis of sodium chloride? The first products of the 
 electrolysis, in this case sodium and chlorine, are called the 
 primary products; any products resulting from a reaction of 
 the primary products with other materials present, in this 
 case, products from the action of sodium on water are called 
 secondary products. Write equations showing the primary and 
 the secondary products of each electrolysis. 
 
 b. To what classes of chemical compounds do conductors of 
 the second class (those which conduct in solution, and at the 
 same time undergo a chemical change) belong? How does the 
 ionic theory explain the passage of the current through the 
 
ROCK SALT 45 
 
 solution? What is the difference between the chloride ion and 
 free chlorine as regards (a) color, (b) conduct when subjected 
 to the influence of a charged plate or pole, (c) chemical prop- 
 erties. What sort of reaction is especially characteristic of 
 those compounds which ionize in solution? 
 
 30. The Purification of Rock Salt, (NaCl). 
 
 Rock salt, as mined, is of course mixed with more or less of 
 those salts which are in solution with it in sea water, for example 
 sulfates and chlorides of calcium, magnesium, potassium, etc. 
 The presence of sulfates in the crude salt may be shown by 
 adding a few drops of barium chloride (BaCl 2 ) solution to a 
 dilute solution of the salt acidified with C. P. hydrochloric acid. 
 A white precipitate indicates sulfates. How may the presence 
 of calcium salts be shown? (Exp. 21.) 
 
 a. Saturate 50 c.c. of water with common salt. From the 
 table of solubilities the amount of salt required may be deter- 
 mined, and a little more than the theoretical amount should 
 be taken to allow for moisture and foreign matter. Filter the 
 solution and test small portions of it for sulfates and calcium. 
 Divide the solution remaining into two equal parts, place one 
 part in an evaporating dish and evaporate on a steam bath 
 until about two-thirds of the water has evaporated, allow the 
 salt to crystallize by cooling, then filter out the crystals, drain 
 as completely as possible, wash them with a very small amount 
 of distilled water, drain, label No. 1, and set away to dry. Into 
 the other half of the solution pass hydrochloric acid gas, gener- 
 ated by the action of the calculated amount of concentrated 
 sulfuric acid (sp. gr. 1.84) diluted with half its volume of 
 water, on 10 grams of sodium chloride. The generating flask 
 should be fitted with a stopper carrying a thistle tube reach- 
 ing to the bottom of the flask, and a delivery tube dipping 
 into the solution of sodium chloride to be purified. Heat the 
 generating flask gently, keeping the stream of gas passing 
 into the salt solution sufficiently rapid so that there is no 
 tendency for the solution to suck back into the generator, and 
 
46 THE HALOGEN FAMILY 
 
 continue until there is no further precipitation in the salt 
 solution. Filter out the precipitate from the solution, drain, 
 wash it with a very small amount of distilled water, and 
 drain, finally allowing it to dry completely by evaporation. 
 Label this sample No. 2. What is this precipitate? Now 
 dissolve small portions of each sample in distilled water 
 and test as before for sulfates and calcium. Judge the rela- 
 tive amounts of impurities in the three samples tested, and 
 draw conclusions as to the efficiency of the two methods of 
 purification. Test solutions of samples 1 and 2 with red and 
 blue litmus paper. What has become of the hydrochloric 
 acid added to sample 2? Is the precipitation of chlorides 
 by hydrochloric acid a general property? Test your conclu- 
 sion by adding 1 c.c. of concentrated hydrochloric acid to 1 
 c.c. of saturated solutions (make these by adding an excess 
 of the solid to 1 c.c. of water, and after shaking five minutes 
 pour off the clear solution from the excess of the solid) of 
 barium chloride (BaCl 2 ) and potassium chloride (KC1). 
 
 31. Potassium Hypochlorite. 
 
 a. Dissolve 5 grams of solid potassium hydroxide (KOH) in 
 15 c.c. of distilled water. Cool the solution and place about 
 half of it in a test-tube, reserving the remainder for Exp. 32. 
 Fit up a small generating flask with a delivery tube and a thistle 
 tube which reaches nearly to the bottom of the flask. Place 
 about 4 grams of finely pulverized manganese dioxide (Mn0 2 ) 
 in the flask and add 10 c.c. of concentrated hydrochloric acid 
 diluted with 10 c.c. of water, through the thistle tube. Warm 
 the flask gently, and when the generation of chlorine begins 
 introduce the delivery tube into the test-tube containing the 
 potassium hydroxide solution. Surround the test-tube with 
 cold water and stir the solution while chlorine is being passed 
 into it. Keep a lively stream of chlorine flowing into the solu- 
 tion for ten to fifteen minutes, then stop the generation of 
 chlorine by allowing the generator to cool down, remove the 
 test-tube and immerse the end of the delivery tube in water or 
 a dilute alkali solution. Do not allow chlorine to escape into the 
 
POTASSIUM CHLORATE 47 
 
 air. This generator may be used again in Exp. 32, but it 
 should not be allowed to stand over night in the desk. 
 
 b. Pour the contents of the test-tube into an evaporating 
 dish and wet a piece of colored cloth in it. Is the cloth bleached? 
 Place the cloth in a beaker containing some dilute sulfuric acid. 
 If no change is observed wet the same piece of cloth with the 
 two solutions again. Explain what happens. Why is sulfuric 
 acid necessary? Explain why this treatment will bleach cloth 
 while air does not. Write equations for the reaction in a, and 
 for the action of sulfuric acid. Why was the test-tube cooled 
 in a ? Acidify a small portion of the solution from a with nitric 
 acid, then add a few drops of silver nitrate solution. What ion 
 does this show to be present? 
 
 32. Potassium Chlorate. 
 
 Using the same generator as in the preceding experiment (31) 
 pass chlorine into the remainder of the potassium hydroxide 
 solution but do not cool the tube. Continue this process until 
 the solution is completely saturated with chlorine, then dis- 
 pose of the generator and contents as in Exp. 27a. How is it 
 possible to tell when the point of saturation is reached? Allow 
 the mixture to cool. What happens? Filter and test the 
 filtrate by adding nitric acid and silver nitrate. What sub- 
 stance is present? Dry the crystals and heat in a dry tube, 
 testing for any gas that might be evolved. What are the 
 crystals? Which is most soluble, potassium chloride or 
 potassium chlorate? Confirm your conclusion by reference 
 to the table. Write all reactions. Explain why this experi- 
 ment and the preceding do not yield the same products. 
 
 33. The Preparation of Potassium Perchlorate. 
 
 When potassium chlorate is heated two reactions take place 
 simultaneously: 4 KC10 3 = 3KC10 4 + KC1 and 2KC10 3 = 2KC1 + 
 30 2 . The greater part of the potassium chlorate will go through 
 the change represented by the first equation if the temperature 
 is held just above the melting-point of the potassium chlorate, 
 so that the evolution of oxygen (represented by the second 
 
48 THE HALOGEN FAMILY 
 
 equation) is very slow. If the temperature becomes too high 
 the evolution of oxygen is rapid, the amount of perchlorate 
 formed is not so great, and it may itself decompose accord- 
 ing to the equation, KC10 4 = KC1 + 20 2 . Thus, in preparing 
 potassium perchlorate from potassium chlorate the potassium 
 chlorate must be heated just enough to keep it melted, and 
 evolving a slow stream of oxygen. 
 
 Place about 8 grams of potassium chlorate (free from any 
 combustible matter, why?) in the hard glass test-tube, and 
 stopper with a cork or rubber stopper carrying a delivery tube 
 leading to the water trough. After testing the apparatus for 
 air tightness melt the potassium chlorate and heat it gently, 
 collecting the oxygen evolved by displacement of water. 
 Take care that the melted salt does not come in contact with 
 the stopper. Why? Regulate the heat so that the evolution 
 of oxygen is very slow, and the salt is kept melted throughout. 
 When 500 c.c. of oxygen have been evolved withdraw the deliv- 
 ery tube and allow the melted mass to cool. From the volume 
 of oxygen evolved calculate the weight of potassium chlorate 
 which has decomposed according to the second equation. 
 Calculate the amount of potassium chloride and potassium per- 
 chlorate which should now qe in the reaction mixture, assuming 
 that all the potassium chlorate has decomposed. Look up 
 the solubilities of the two salts left and devise a method for 
 separating these. After submitting your method for the ap- 
 proval of an instructor carry out the separation, wash the po- 
 tassium perchlorate and weigh roughly. What per cent of the 
 theoretical yield is obtained? Should the method of sepa- 
 ration used give a pure perchlorate? Test a small portion 
 for chlorides and save the rest of the product for Exp. 34. 
 
 34. To Test the Purity of the Potassium Perchlorate. Quanti- 
 tative. 
 
 Place from 0.2 to 0.5 gram of the perfectly dry potassium 
 perchlorate obtained in Exp. 33 in a weighed hard glass ignition 
 tube, weigh accurately, then heat until the evolution of oxygen 
 
BROMINE 
 
 49 
 
 is complete, and weigh the tube and residue. Calculate the 
 per cent of oxygen in the given sample. Compare this result 
 with that obtained in Exp. 24 and with the theoretical compo- 
 sition of KC10 4 . Was the potassium perchlorate pure? 
 
 35. The Preparation of Bromine. 
 
 Caution: Do not breathe the vapor of bromine. Do not 
 get liquid bromine on the skin, as it causes severe burns. 
 
 Calculate the quantities of manganese dioxide (Mn0 2 ), 
 potassium bromide (KBr) and sulfuric acid necessary to 
 make 2 c.c. of liquid bromine (sp. gr. 3.18). Have this calcula- 
 
 FIG. 14. 
 
 tion approved by an instructor before proceeding. The 
 volume of concentrated sulfuric acid required (sp. gr. 1.84) 
 should be diluted with one-half its volume of water by pouring 
 the measured amount of concentrated acid into the measured 
 amount of water in a beaker. Never pour water into sulfuric 
 
50 THE HALOGEN FAMILY 
 
 acid. Grind up the required amount of potassium bromide in 
 the mortar and mix with the manganese dioxide, then place the 
 mixture in a retort (storeroom) and add the diluted sulfuric 
 acid through the tubule. Close the tubule and heat gently, 
 protecting the bottom of the retort by a wire gauze. Collect 
 the bromine in a flask into which the neck of the retort extends 
 as far as possible, and which is kept cool by being placed under 
 a stream of running water, while the escape of bromine vapor 
 into the air is prevented by folding a piece of wet filter paper 
 over the mouth of the receiving flask (Fig. 14). When the evo- 
 lution of bromine has ceased remove the flask and test the 
 gas in the flask with pieces of paper wet with starch solution, 
 and with starch potassium iodide solution (see Exp. 25) . (Then 
 proceed at once to Exp. 36. Do not allow the bromine to stand 
 over night.) To what is the difference in color in the two tests 
 due? Refer to the starch potassium iodide test for chlorine. 
 Would you expect to find the same succession of colors in this 
 case? Write equations for the reactions involved. Could 
 bromine be prepared by any of the methods used for chlorine? 
 If potassium iodide was present as an impurity in potassium 
 bromide how could the potassium bromide be most readily 
 purified? Is there any connection between the low boiling- 
 point of bromine and the fact that the air above the liquid 
 bromine is strongly colored with bromine vapor at room 
 temperature? 
 
 36. Potassium Bromate. 
 
 a. Cover the bromine collected in the preceding experiment 
 with 2 to 4 c.c. of water and add drop by drop a strong solution 
 of potassium hydroxide (KOH, 1:1) with shaking, until the 
 liquid bromine has entirely dissolved, and the color of the solu- 
 tion is only faintly yellow. What reaction is taking place? 
 (Note: If a brown precipitate of manganese dioxide, carried over 
 mechanically by the bromine, appears in the solution on the 
 addition of potassium hydroxide it must be filtered out before 
 the color of the bromine in solution is entirely removed by the 
 
POTASSIUM BROMATE 51 
 
 potassium hydroxide.) Heat this solution to boiling, (what 
 further reaction occurs?), and continue boiling until crystals 
 begin to separate out, then cool under running water, and filter 
 after the solution has come to room temperature. Rinse out 
 the flask and the precipitate on the filter paper with 2 c.c. water 
 and save the filtrate. Allow the crystals to drain as com- 
 pletely as possible and then in order to purify them wash 
 into an evaporating dish with a fine stream of distilled 
 water from the wash bottle. Heat till the crystals are 
 entirely dissolved, then allow the solution to crystallize 
 spontaneously by cooling and evaporation. Remove the 
 crystals from the solution before all the water has evapo- 
 rated and dry them on filter paper, then place in a weighed 
 (quantitative) test-tube and label in such a way that the label 
 can be readily removed before a second weighing. Boil the 
 filtrate from the first crystallization (in an evaporating dish) 
 until it begins to crystallize, then allow to cool. Remove the 
 crystals, dry them on filter paper, and place them in another 
 weighed test-tube, labeling as before. These may be best 
 labeled according to solubility. Which crystals represent the 
 least soluble, and which the most soluble portion? Write 
 the reaction involved when bromine acts on potassium hydrox- 
 ide under the given conditions, and after looking up the solu- 
 bilities of the products determine which salt you will expect 
 to find in the least, and which in the most soluble portion of 
 your product. Weigh the test-tubes with the salt, and by 
 subtracting the weight of the empty tubes get the weights of 
 the salts. 
 
 b. Devise some method of testing these two salts to determine 
 whether your conclusions regarding their identity are correct. 
 Get the approval of an instructor before proceeding to apply the 
 tests you devise. Compare the reaction between bromine and 
 potassium hydroxide with that used in Exp. 32 between 
 chlorine and potassium hydroxide. Calculate from the original 
 amount of KBr taken in Exp. 35, the theoretical yields of 
 KBr0 3 and KBr. What per cent of the theoretical yield is 
 
52 THE HALOGEN FAMILY 
 
 the amount obtained; or in other words, what is your per- 
 centage yield? At what stages of the process has there been an 
 opportunity for loss of material? 
 
 37. Iodine. 
 
 Mix together a small crystal of potassium iodide, and a 
 little manganese dioxide. Place the mixture in a test-tube 
 and add a few drops of concentrated sulfuric acid, then heat 
 gently. Note the color of the gas evolved. Is it lighter or 
 heavier than air? What observation leads to this conclusion? 
 Hold a paper wet with starch solution in the gas. Explain. 
 Why was potassium iodide used with the starch solution in the 
 test for chlorine and bromine and not in this test? Compare 
 this method for the preparation of iodine with those for chlorine 
 and bromine. Why would it not be safe to use potassium 
 chlorate as the oxidizing agent on a mixture of potassium 
 iodide and concentrated sulfuric acid (see test-book under 
 chlorine dioxide). What would be the reaction between KC10 3 
 and HI? 
 
 38. Hydrofluoric Acid and its Etching Action on Glass. 
 
 Prepare hydrofluoric acid by placing 2 grams of calcium 
 fluoride (CaF 2 ) in a lead dish, or evaporating dish which has 
 been coated with paraffin, and mixing with enough concentrated 
 sulfuric acid to make a paste. Cover the dish with a glass 
 plate which has been protected by a coating of paraffin except 
 in the places which are to be etched. Allow the dish and 
 cover to stand for about an hour, then remove the paraffin 
 from the cover glass and examine. Write the reactions and 
 account for the disappearance of part of the glass. Test the 
 gas in the dish by breathing across it. What are the fumes? 
 Hold a rod wet with ammonium hydroxide (NH 4 OH) above 
 the dish. What are the fumes? Is this a distinctive test for 
 hydrofluoric acid? For what other acid has it been used? 
 What one property is essential that an acid should give this 
 test? 
 
EQUILIBRIUM 53 
 
 39. Equilibrium. 
 
 a. Review the preparation of hydrochloric acid and show how 
 the reaction was proved to be reversible. What property of 
 the hydrochloric acid made it possible to get complete con- 
 version of the sodium chloride into hydrochloric acid, even 
 though the reaction was reversible? 
 
 The reaction, NaH 2 P0 4 + HC1 NaCl + H 3 P0 4 , is revers- 
 sible, and if at any fixed temperature solutions of primary 
 sodium phosphate and hydrochloric acid, or of sodium chloride 
 and phosphoric acid of equivalent concentrations are mixed, 
 the reaction mixture will consist of a solution of the two 
 original materials and the two resulting products in exactly 
 the same proportions, regardless of which pair of materials 
 was originally used. 
 
 b. To show this experimentally make an indicator by mixing 
 six volumes of potassium iodide solution with one volume of 
 potassium bromate solution, and add a few drops of this indi- 
 cator to small portions of normal hydrochloric and phosphoric 
 acid solutions, noting the effect of the indicator on the two 
 acids. Now prepare four tubes, in one placing equal quantities 
 of normal solutions of primary sodium phosphate and hydro- 
 chloric acid, in another the same quantities of normal solutions 
 of sodium chloride and phosphoric acid, in the third the same 
 quantities of normal hydrochloric acid and water, in the fourth 
 the same quantities of normal phosphoric acid and water, and 
 label each tube. Shake each thoroughly then add an equal 
 number of drops 1 of the indicator to each (the two diluted 
 acid solutions are to be used as standards to show the color 
 given by hydrochloric and phosphoric acids of the same con- 
 centration as in the first two tubes) . From the color reactions 
 (these must be noted at once, as the color deepens on standing) 
 draw conclusions as to the composition of the solution in each 
 of the first two tubes. 
 
 student should provide himself with half a dozen pieces of glass 
 tubing, drawn to a jet at one end and kept clean and dry to use as drop- 
 ping tubes. 
 
54 THE HALOGEN FAMILY 
 
 c. Add twice normal phosphoric acid to an equal volume of 
 normal sodium chloride and note the color given by the same 
 number of drops of the indicator as before. Compare with the 
 tube made from normal phosphoric acid and sodium chloride. 
 What effect does increased concentration of one of the reacting 
 substances have on the point of equilibrium? 
 
CHAPTER VII 
 SULFUR 
 
 40. Physical Properties of Sulfur. 
 
 What is a polymorphous substance? Do the polymorphic 
 forms of a given substance differ most with respect to physical 
 or chemical properties? Illustrate. Place about 5 grams 
 of sulfur (sample A) in a short, wide, dry test-tube and heat 
 gently till it just melts (use special care not to heat the sulfur 
 much above its melting-point), then pour a small portion of 
 the melted sulfur into a beaker containing cold water (sample 
 B) . Continue the heating of the rest of the sulfur in the test- 
 tube until it begins to boil, noting the changes in appearance 
 and viscosity as the temperature rises. Pour the boiling 
 sulfur into another beaker of cold water (sample C). Com- 
 pare these three samples of sulfur (A, B, C) as to appearance, 
 texture and solubility in carbon disulfide. To study this 
 last property place a small portion of each sample in separate 
 test-tubes and add 1-2 c.c. of carbon disulfide (CS 2 ), then 
 shake, taking care that the test-tubes are at a safe distance 
 from all flame, as carbon disulfide is very inflammable, and 
 its vapor mixed with air is explosive. Allow small portions of 
 these solutions to evaporate spontaneously. Did any sulfur 
 dissolve? Record the results, then pour the contents of the 
 test-tubes down the sink, following with plenty of water. Let 
 portions of samples B and C stand in the desk over night and 
 examine again in the same way. Explain the differences in 
 the behavior of the three samples. Retain the test-tube in 
 which the sulfur was melted for use in the next experiment 
 (41 6.) 
 
 55 
 
56 SULFUR 
 
 41. Chemical Properties of Sulfur and Hydrogen Sulfide. 
 
 a. Mix 1 gram powdered zinc and 0.5 gram sulfur, place on a 
 crucible cover and heat. Caution. What is the reaction? 
 
 b. Mix 2 grams of iron filings and 1 gram of sulfur, place in 
 the test-tube used in Exp. 40 and heat. What is the reaction? 
 
 c. What reactions of chlorine and oxygen are analogous to 
 these for sulfur? From the position of sulfur in the periodic 
 table would you expect it to be like these two elements? 
 Examine these two products by placing samples of each (obtain 
 the iron compound by breaking the tube) in separate test-tubes 
 and adding a few c.c. of dilute hydrochloric acid. Note the odor 
 of the gas evolved (cautiously). Test it with paper moistened 
 with lead acetate solution. Write the reactions involved. 
 
 d. Introduce a small portion of burning sulfur on the iron 
 spoon into a gas-collecting bottle, and cover the bottle as com- 
 pletely as possible with a glass plate. What is the reaction? 
 When the reaction has ceased place 10 c.c. of water in the bottle 
 and shake. What is formed? Pour this solution into a test- 
 tube and pass in hydrogen sulfide, which may be generated by 
 placing some of the iron or zinc sulfide in a test-tube, adding 
 dilute hydrochloric acid and closing the tube with a stopper 
 carrying a delivery tube which can be inserted into the solution. 
 Note the appearance and odor of the solution after the reaction 
 has taken place. Write the equation. 
 
 e. This is an example of a colloidal solution, and is called 
 "milk of sulfur." Two of the general properties of colloidal 
 solutions may be easily illustrated by pouring the solution 
 through a filter paper, showing the fine state of division of the 
 sulfur, then adding a few crystals of sodium chloride, shaking 
 and filtering again. The electrolyte has coagulated the colloid. 
 
 /. Pass some hydrogen sulfide, generated in the same way, 
 into 10 c.c. of water, stopper and allow the solution to stand 
 over night. What is formed? What is the oxidizing agent in 
 this case? These reactions have been suggested to explain the 
 occurrence of free sulfur in volcanic regions. 
 
 g. Pass hydrogen sulfide from a Kipp generator or from the 
 
SULFUR DIOXIDE 57 
 
 laboratory supply into a solution of a zinc salt, for example 
 zinc chloride. Is the reaction reversible? 
 
 42. Sulfur Dioxide. 
 
 a. If fresh solutions of ferric chloride (FeCl 3 ) and potassium 
 ferricyanide (K 3 Fe(CN) 6 ) are mixed, and the brown solution 
 so obtained is taken up on filter paper, it may be used as an 
 indicator for sulfur dioxide (S0 2 ), as the color changes to a 
 bluish-green when exposed to the gas. The color changes in the 
 same way if the paper is exposed to air and light for any length 
 of time, so it must be prepared and used for the test at once. 
 If this is done there is no difficulty in recognizing appreciable 
 amounts of sulfur dioxide. 
 
 b. Heat 2 c.c. of concentrated sulfuric acid in contact with 
 copper turnings in a test-tube until there is a spontaneous 
 evolution of gas. Test the escaping gas with the freshly 
 prepared paper. What is formed? Write the reaction. Does 
 dilute sulfuric acid act on copper? (See Exp. 13.) Allow the 
 tube and contents to cool, then pour the contents in a beaker 
 containing about 50 c.c. of water. Allow the precipitate to 
 settle, and explain the color of the solution. 
 
 c. Add concentrated sulfuric acid to a small crystal of potas- 
 sium bromide (KBr) and heat gently, testing the escaping gas as 
 before. Explain the color of the evolved gas in this case. 
 Write the reactions taking place. 
 
 d. What name would be given to any process by which S0 2 is 
 obtained from H 2 S0 4 ? In these two experiments what are the 
 reducing agents? What is the oxidizing agent? 
 
 e. Add some dilute hydrochloric acid to a few sodium sulfite 
 (Na 2 S0 3 ) or sodium acid sulfite (NaHS0 3 ) crystals in a test- 
 tube, warm and test for sulfur dioxide. Write the reaction. 
 Fit this tube with a cork and delivery tube, and pass the gas 
 evolved (use care that none of the liquid is carried over) into 
 a test-tube containing 9 c.c. of water and 1 c.c. ammonium 
 hydroxide (NH 4 OH) until the solution becomes saturated. 
 Now pass the gas into another test-tube containing 1 c.c. 
 
58 SULFUR 
 
 potassium permanganate (KMn0 4 ) solution and 9 c.c. water. 
 Note the change in appearance. (What other reagent has 
 the same effect on KMn0 4 ?) 
 
 Compare the two solutions by adding to each a little barium 
 chloride (BaCl 2 ) solution, shake and note the effect, then allow 
 to settle, decant the clear liquid and add dilute C.P. hydro- 
 chloric acid to the precipitate in each tube. How do you 
 account for the difference in behavior? Write equations for 
 all the reactions involved. 
 
 43. Preparation and Properties of Sodium Thiosulfate. 
 
 Dissolve 10 grams of anhydrous sodium carbonate in 50 c.c. 
 of water, warming if necessary, or if only the hydrated salt 
 (Na 2 C0 3 -10 H 2 0) is available dissolve a quantity of it equivalent 
 to 10 grams of Na 2 C0 3 in a correspondingly smaller amount of 
 water. Divide this solution into two equal parts, and saturate 
 one part with sulfur dioxide, generated by heating 10 grams 
 of copper turnings with 20 c.c. concentrated sulfuric acid in a 
 generating flask provided with a thistle tube and delivery tube. 
 Be particularly careful to see that the solution of the carbonate is 
 not at any time sucked back into the generating flask. Add to 
 the solution prepared in this way (what is it?) the other part 
 of the sodium carbonate solution, and heat to boiling. What 
 two reactions have occurred thus far? Write equations for 
 each. After the effervescence has ceased (to what is it due ?) 
 add 4 grams of sulfur, boil for 20 minutes keeping the volume 
 of the solution constant, filter off the excess of sulfur, evapo- 
 rate the solution over a water bath until crystallization com- 
 mences, then allow it to cool. Filter off the crystals and evapo- 
 rate the filtrate to crystallization a second time, then cool and 
 combine the two crops of crystals. Weigh the product (after 
 drying) on the rough laboratory balance, and calculate the 
 percentage yield. Dissolve some of the crystals in water, 
 and test the solution with blue and red litmus paper. 
 Hydrolysis of the salt into a base and acid has taken place. 
 According to the test, which (base or acid) is more highly 
 
SULFURIC ACID 59 
 
 ionized? Add a few drops of hydrochloric acid to the solution 
 and warm. The precipitate is milk of sulfur, where have you 
 seen this before? Write the equation for the reaction. Is 
 thiosulfuric acid a stable compound? Add a solution of so- 
 dium thiosulfate to some freshly prepared silver chloride and to 
 a small crystal of iodine. Write the reactions. Of what tech- 
 nical use is sodium thiosulfate? 
 
 44. Sulfuric Acid. 
 
 a. Study the properties of sulfuric acid by the following tests : 
 Dip a strip of filter paper into dilute sulfuric acid and dry 
 on the radiator, then examine. Add a few drops of concen- 
 trated sulfuric acid to a little sugar in an evaporating dish. 
 Compare with the effect of heating sugar on a crucible lid. 
 Add 1 c.c. of concentrated sulfuric acid to 1 c.c. of water in 
 a test-tube, noting the temperature before and after mixing. 
 Does the last experiment throw any light upon the nature of 
 the change in the first two? Explain. 
 
 b. Add barium chloride (BaCl 2 ) solution to dilute sulfuric 
 acid. What is the reaction? Allow the precipitate to settle, 
 pour off the clear liquid and add a few cubic centimeters of con- 
 centrated C. P. hydrochloric acid to the precipitate. Is the 
 above reaction appreciably reversible under these conditions? 
 
 45. Neutralization. Determination of the Concentration of a 
 Solution of Acid or Base. Quantitative. 
 
 The neutralization of an acid by a base depends on the union 
 of the hydrogen ions of the acid with the hydroxyl ions of the 
 base to form unionized water. If an indicator, litmus, is 
 present in a solution of an acid it will give the solution a red 
 color, due to the H + ions. If a base, as KOH, is added to that 
 acid solution, the H + ions are gradually removed by the OH" 
 ions of the base, until finally there are none left in the solution, 
 and the addition of the next drop of the solution of the base 
 causes the litmus to turn blue, because of the free OH~ ions 
 in the solution. Such a solution, the addition of one drop of 
 
60 SULFUR 
 
 base or acid to which causes it to turn litmus blue or red, is 
 said to be neutral, and is free from both H + and OH~ ions 
 except to the extent to which they occur in water, where they 
 are equal in number. What ions are present in a solution of 
 H 2 S0 4 which has been neutralized by KOH? Write the ionic 
 equation for the reaction. 
 
 The concentration of an acid solution may be determined 
 by neutralizing a measured volume of the solution, evaporating 
 to dryness and weighing the salt formed (it must be a non- 
 volatile, stable salt, why?) and calculating from the weight of 
 the salt, the weight of the acid originally in the solution. 
 
 a. Make a dilute solution of sulfuric acid by pouring 3 c.c. 
 of concentrated sulfuric acid into 100 c.c. of water and mix 
 thoroughly by shaking. Place in a flask and stopper the flask, 
 so that the solution will not change in concentration by 
 evaporation. Dissolve about 6 grams of solid potassium 
 hydroxide (KOH) (use care in handling this, as it is a caustic 
 and burns the skin) in 100 c.c. of water, shake thoroughly 
 and stopper. Label the flasks. 
 
 Save these solutions as they must be used again in Exp. 46. 
 
 b. Rinse out and fill a burette (storeroom) with the prepared 
 potassium hydroxide solution drawing enough solution out 
 through the nozzle to insure the removal of all the air from the 
 fine tubing, and bring the solution in the burette to the zero 
 mark. Now make a rough determination of the amount of 
 potassium hydroxide solution required to neutralize 10 c.c. of 
 the sulfuric acid solution as follows: Measure out 10 c.c. of 
 the sulfuric acid solution with the pipette, and place it in a 
 beaker, add 2 drops of phenolphthalein (which is an indicator 
 like litmus, except that it is colorless in acid and red in alkaline 
 solutions) and then while stirring continuously, add the 
 potassium hydroxide solution from the burette, 1 c.c. at a time, 
 until the solution in the beaker becomes permanently red, i.e., 
 until it is alkaline. Read the burette. The volume of 
 potassium hydroxide solution used is not more than 1 c.c. 
 greater than the volume required for exact neutralization of 
 
NEUTRALIZATION 61 
 
 10 c.c. of the sulfuric acid solution if these directions have 
 been carefully followed. Now make the exact, final deter- 
 mination as follows: 
 
 c. Clean and weigh (quantitative) an evaporating dish and 
 cover, then place in it exactly 10 c.c. of the sulfuric acid 
 solution, measured with the pipette, and weigh again. From 
 the weight of the solution and its volume calculate its specific 
 gravity. Add a couple of drops of phenolphthalein and after 
 reading the burette neutralize the acid with potassium 
 hydroxide from the burette, adding at once 1 c.c. less than the 
 amount required in the rough neutralization, then to determine 
 the fraction of a cubic centimeter required to just neutralize 
 the acid add the potassium hydroxide drop by drop, stirring 
 constantly with a glass rod, and stop the addition the moment 
 the solution turns permanently red. The color is best seen 
 against a white background. 
 
 When the solution is just neutral read the burette again, and 
 determine the volume of potassium hydroxide solution used. 
 Evaporate the neutral solution to dryness on a steam bath, or 
 by setting it above a beaker of boiling water, and when dry 
 heat the covered dish gently with the direct flame, cool and 
 weigh. Heat the dish and contents again, and weigh, until 
 two successive weighings give the same result. What is left 
 in the dish? 
 
 d. Define per cent composition and normal solution. From the 
 weight of the salt formed calculate the weights of sulfuric 
 acid and potassium hydroxide used. Calculate the normal 
 strength of the sulfuric acid and potassium hydroxide 
 solutions. How many milligrams of sodium hydroxide (NaOH) 
 will 1 c.c. of the sulfuric acid neutralize? How many milli- 
 grams of ammonium hydroxide (NH 4 OH)? How many milli- 
 grams of nitric acid (HN0 3 ) will 1 c.c. of the potassium 
 hydroxide neutralize? Of hydrochloric acid? Of sulfurous 
 acid (H 2 S0 3 )? Calculate the per cent composition of the 
 sulfuric acid solution. What further information would be 
 necessary for you to calculate the per cent composition of the 
 
62 SULFUR 
 
 potassium hydroxide solution? Why could not the concen- 
 tration of these solutions be determined by evaporating off 
 the water and weighing the residue, as in Exp. 18? 
 
 46. Potassium Acid Sulfate. 
 
 Place 10 c.c. of the dilute sulfuric acid prepared in Exp. 45 
 (measured with the pipette) in a clean evaporating dish, then 
 add, from the burette, just half as much potassium hydroxide 
 solution as was found necessary for neutralization, and evapo- 
 rate the solution to dryness on the steam bath. What is 
 formed? Dissolve a few of the crystals in water and test the 
 solution with red and blue litmus paper. Explain the effect. 
 Dip a filter paper into the solution, and allow it to dry out on 
 the radiator. Is sulfuric acid present in the solution? If 
 sulfuric acid is not found, how can the effect on litmus be 
 explained? Add some of the solution to some solid sodium 
 carbonate in a test-tube. Explain. Add barium chloride 
 (Bad 2 ) to the solution, then hydrochloric acid. Is the sulfate 
 ion present? 
 
CHAPTER VIII 
 
 NITROGEN 
 
 47. Preparation and Properties of Nitrogen and Ammonia. 
 
 a. What are the principal constituents of the air? Are 
 they present as a mixture or in chemical combination? 
 How may oxygen be removed from the air? 
 
 b. Float a crucible on the trough full of water, place in it about 
 0.5 gram red phosphorus and light it by touching it with the 
 hot end of a file or iron spoon. Quickly invert a gas-collecting 
 bottle over the dish, bringing the mouth of the bottle below the 
 surface of the water, and hold in that position till the phosphorus 
 has ceased burning, and the water inside the bottle has come to a 
 fixed level. Account for the rise of the water in the bottle. 
 Place a glass cover over the mouth of the bottle while it is still 
 under water, and withdraw. Test the gas in the bottle by in- 
 serting into it a burning splinter. Does it support combustion? 
 Does it burn? Insert into it a piece of burning magnesium 
 ribbon. Does it support combustion in this case? What is 
 formed? 
 
 c. Place about 0.5 gram of magnesium powder in a crucible, 
 cover and heat while covered until all glowing has ceased. 
 The crucible contains a mixture of what two compounds? 
 Place the contents of the crucible in a test-tube, moisten with 
 water, and warm gently. Test the escaping gas with moistened 
 red litmus paper. What is it? Write the reactions. 
 
 d. To 1 c.c. of ammonium chloride (NH 4 C1) solution in a test- 
 tube add a few drops of potassium or sodium hydroxide solution, 
 and warm, testing the gas evolved as before. What is it? 
 Write the reaction. Hold a rod wet with concentrated hydro- 
 chloric acid at the mouth of the test-tube. Explain the fumes 
 formed. Devise and apply a test to determine whether the 
 gas is lighter or heavier than air. 
 
 63 
 
64 NITROGEN 
 
 48. The Weight of a Liter of Ammonia. Quantitative. 
 
 Generate ammonia by warming together in flask fitted with a 
 delivery tube, and proved to be air tight, a mixture of equal 
 weights of ammonium chloride and powdered quick lime (CaO). 
 Collect three test-tubes of the gas by air displacement and test 
 its solubility in water, dilute alkalis (dilute sodium hydroxide 
 or potassium hydroxide) and dilute acids. From this test 
 determine which will make the most suitable liquid for absorbing 
 the gas, then proceed with the determination of the weight of a 
 liter, according to the method used for chlorine. Could con- 
 centrated sulfuric acid be used to dry this gas? Why? If the 
 gas is prepared as directed, from dry ammonium chloride and 
 quicklime it will not be necessary to dry it, but it may be passed 
 directly into the weighed flask. Should the gas be collected 
 with the flask held upright or inverted? The flask is sufficiently 
 filled when a strip of moistened red litmus paper held at the 
 mouth turns blue. 
 
 49. Ammonium Hydroxide. 
 
 a. Attach a U tube containing water to the ammonia generator 
 used in the previous experiment, similar to the arrangement used 
 for making a solution of hydrochloric acid (Fig. 13) and pass in 
 ammonia until the solution is saturated. Note the effect on 
 the temperature of the solution. Test this solution with litmus, 
 then boil a small portion in a beaker for 5 minutes and test it 
 again. Explain. Neutralize another portion of the solution 
 with dilute hydrochloric acid and evaporate the neutral solution 
 to dryness on the water bath, then heat the dry residue with the 
 direct flame. Repeat this procedure only neturalize with dilute 
 sulfuric acid instead of hydrochloric. Account for the difference 
 in behavior on heating the dry salts. Is there any relation be- 
 tween the volatility of the acid and of its ammonium salt? 
 Allow a small sample of ammonium carbonate to stand on a 
 watch glass over one laboratory period, and account for the 
 change. 
 
 b. Add ammonium hydroxide solution one drop at a time to 
 
NITRIC ACID 65 
 
 freshly precipitated silver chloride and to copper sulfate solution. 
 These reactions are probably due to the presence of ammonia, 
 (NH 3 ) in the solution, and the soluble compounds formed are 
 AgCl-2NH 3 and Cu(NH 3 ) 4 S0 4 . 
 
 50. Nitric Acid and the Oxides of Nitrogen. 
 
 a. Mix 10 grams of sodium nitrate (NaN0 3 ) with about 8 c.c. 
 of concentrated sulfuric acid in a retort, arrange a. receiving 
 flask which can be cooled by running water, as in Exp. 35, and 
 warm gently, protecting the retort with a wire gauze, until the 
 nitric acid has been entirely distilled over. This should give 
 about 5 c.c. of concentrated nitric acid. Note the color of 
 the acid and of the gas above it. Pure nitric acid is colorless. 
 Can you explain the color here? (If this acid cannot be used 
 at once dilute it with one-half its volume of water before putting 
 it away.) 
 
 b. Place about 1 c.c. of this acid in a test-tube, and dilute it 
 with one-half its volume of distilled water. Moisten a piece 
 of filter paper with this solution, then allow it to dry out on 
 the radiator. Dilute a few drops of the acid with 2-3 c.c. of 
 water and add barium chloride (BaCl 2 ) solution. Now add a 
 small amount of sulfur to a new portion of the diluted acid, warm 
 till reaction seems to cease, noting the appearance of the solution 
 and of the gas evolved, and again test the solution with filter 
 paper as before (see Exp. 44). Account for the difference in the 
 two tests. What has become of the sulfur? Add a few drops 
 of barium chloride solution to the mixture in the test-tube 
 (after diluting with distilled water) and explain the change. 
 Does nitric acid form a precipitate with barium chloride? 
 What property of nitric acid does this reaction with sulfur 
 illustrate? 
 
 c. Nitric Oxide. Arrange a generating flask as usual, with 
 thistle tube and delivery tube. Place about 3 grams of copper 
 in the flask and add the rest of the nitric acid prepared above, 
 diluted with one-half its volume of water. Heat the flask and 
 contents and collect one bottle of the gas evolved by displace- 
 
66 NITROGEN 
 
 ment of water, noting the color of the gas in the generating 
 flask, and in the gas collecting bottle. Account for the differ- 
 ence in color. At the same time collect a test-tube one-third 
 full of the gas, note its color, and while the mouth of the tube is 
 still under water pour into the tube (by pouring upward) the 
 air from another test-tube which is two-thirds full of air, the 
 rest of the tube being full of water and inverted in the bath 
 like the first one. Account for the change in appearance and 
 volumes of the gases after mixing. Shake the tube while still 
 inverted under water to determine whether the gas formed is 
 soluble in water. Does this experiment throw any light on the 
 reason for the difference in cclor between the gas in the generat- 
 ing flask and in the collecting bottle? Write the reactions 
 involved. Will the nitric oxide support the combustion of a 
 burning splinter? Of phosphorus? (Test both of these on the 
 bottle of gas.) 
 
 d. Nitrous Oxide. Fit a small flask with stopper and delivery 
 tube leading to the water trough. Place about 10 grams of 
 ammonium nitrate (NH 4 N0 3 ) in the flask and heat gently 
 (too rapid heating may cause an explosion) collecting two 
 bottles of the gas evolved by displacement of water. Note 
 the color of the gas. Does it support the combustion of 
 wood and phosphorus? Compare it with oxygen in this 
 respect. Is there any apparent reaction with the oxygen of 
 the air? Compare it with nitric oxide in this respect. How 
 could a given sample of gas be tested to determine whether it 
 was nitric oxide or nitrous oxide? 
 
 61. Behavior of Nitrates and Nitrites on Heating. 
 
 Heat samples of sodium nitrate, sodium nitrite and powdered 
 lead nitrate in hard glass test-tubes. Note the color of the gas 
 evolved and test it for oxygen. Add 1 c.c. of dilute hydrochloric 
 acid to the residue after heating, in each case, again noting the 
 color of the gas evolved. Explain the reactions and write 
 equations. Arrange in tabular form the products obtained on 
 heating sodium nitrate, ammonium nitrate, ammonium nitrite, 
 
OXIDATION WITH NITRIC ACID 67 
 
 sodium nitrite and lead nitrate, and the products obtained when 
 the residue from heating is treated with dilute hydrochloric 
 acid. 
 
 The Combining Weight of a Metal by Oxidation with Nitric 
 Acid. 
 
 Combining Weight of Tin. Quantitative. Carefully clean and 
 dry an evaporating dish and cover glass, weigh it, then place 
 in it about 1 gram of tin foil and weigh again. Convert the 
 metal to its oxide by adding dilute nitric acid (HN0 3 ), about 
 5 c.c. at a time, until all reaction apparently ceases. (Note 
 the color of the gas evolved and explain.) Then remove the 
 cover and evaporate on a steam bath to remove the excess 
 of nitric acid and water, finally heating the covered evaporating 
 dish with the direct flame to remove the last traces of water. 
 Weigh, heat again and reweigh until the dish and contents 
 have come to constant weight. The product in the dish is an 
 oxide of tin. From the data obtained calculate the combining 
 weight of the tin. If any nitrate of tin was formed during 
 the process what has become of it? Was there any indication 
 that a nitrate was being decomposed while the dish was heated 
 with the direct flame? Write the reactions involved. The 
 formula for the tin oxide may be determined from the combin- 
 ing proportions and the atomic weights of tin and oxygen. 
 
CHAPTER IX 
 THE ATMOSPHERE 
 
 53. The Determination of Oxygen in the Air. Quantitative. 
 
 If a eudiometer is not available for this experiment a test-tube 
 may be roughly graduated in 10 c.c v marking it with a pencil 
 for writing on glass, or with file marks, and used in its place. 
 
 v 
 
 FIG. 15. 
 
 Invert a eudiometer containing the air of the laboratory over a 
 \vater bath, clamp it in position, read the volume of air at at- 
 mospheric pressure, and insert into the air a bit of yellow 
 phosphorus fastened on the end of a wire. Fig. 15. The 
 phosphorus may be prepared by melting a piece of yellow 
 phosphorus about the size of a pea, under water, placing the 
 wire in the melted phosphorus and allowing the phosphorus to 
 solidify in this position. Yellow phosphorus should be kept under 
 water and should not be handled with the bare hands, as it causes 
 very unpleasant burns. Allow this to stand until there is no 
 further diminution in volume of the gas and again read the volume 
 
 68 
 
COMPOSITION OF AIR 69 
 
 at atmospheric pressure. Will the gas remaining in the tube 
 support combustion? What is it? What is the per cent by 
 volume of oxygen in the air? In this experiment the air was 
 always saturated with moisture, and was measured at at- 
 mospheric temperature and pressure. Is it necessary to calcu- 
 late the volume of the dry gas under standard conditions, before 
 and after absorption of the oxygen, before calculating its per 
 cent by volume? Would it be necessary to do so in order to 
 calculate its per cent by weight? 
 
 54. The Composition of Air Dissolved in Water. Quantitative. 
 
 The composition of air dissolved in water may be very easily 
 determined by the method used for the determination of oxygen 
 in atmospheric air, the only change in manipulation being the 
 source of the sample for analysis. Use the apparatus shown in 
 Fig. 16, in which the eudiometer or graduated test-tube is fitted 
 with a two hole stopper, one hole of which carries a tube leading 
 just to the lower surface of a cork fitted to a 500 c.c. Erlenmeyer 
 flask, while the other hole is fitted with a tube bent so that it 
 leads under the surface of water in a beaker. Each of these con- 
 necting tubes must be in two pieces, connected with a rubber tube 
 which can be closed with a pinch cock. Fill the whole apparatus 
 completely with fresh tap water and clamp the eudiometer in 
 an upright, inverted position above the Erlenmeyer flask, open 
 both the pinch cocks (if water runs out on opening the cocks it 
 is an indication of a leak, which must be corrected before 
 proceeding) and slowly heat the water in the Erlenmeyer flask. 
 The gas which separates will rise to the top of the eudiometer, 
 while a small amount of water will run out into the beaker be- 
 cause of the expansion of the water on heating and its dis- 
 placement by the gas. The water may be brought to the boiling- 
 point, but it must not be allowedlo boil lest steam escape, carry- 
 ing with it some of the gases. If this does not supply a sufficient 
 sample of gas for analysis the two pinch cocks may be closed, 
 the Erlenmeyer flask refilled with fresh tap water, taking care 
 that it is completely filled with water when the eudiometer is 
 
70 
 
 THE ATMOSPHERE 
 
 replaced, and the procedure repeated until a sufficient sample is 
 obtained. When a sufficient sample of the gas is obtained (20 
 c.c. will be enough for rough results) close both the pinch cocks, 
 and bring the eudiometer, with its stopper and connecting tubes, 
 under the surface of water in a water bath, then remove the 
 stopper and read the volume after the gas has come to at- 
 mospheric temperature, and proceed to determine the oxygen as 
 
 FIG. 16. 
 
 in Exp. 53. Calculate the per cent composition by volume of 
 the gas analyzed, and compare with the per cent composition 
 calculated from the solubility of the two gases and their partial 
 pressures in the air. What would be the relative amounts 
 dissolved if the air were 40 per cent oxygen? State Henry's 
 law. 
 
CHAPTER X 
 PHOSPHORUS, ARSENIC, ANTIMONY AND BISMUTH 
 
 55. The Oxides of Phosphorus. 
 
 Connect a hard glass combustion tube open at each end with a 
 delivery tube fitted with a pinch cock and leading to the bottom 
 of a test-tube or small flask one-third full of water which is fitted 
 with a two hole stopper, the other hole of which carries a tube 
 connected with a suction pump or aspirator bottle (Fig. 17). 
 
 FIG. 17. 
 
 a. Place red phosphorus the size of a pea in the hard glass 
 tube, light by touching with a red hot file, and by manipulating 
 the pinch cock keep the phosphorus burning quietly while air is 
 drawn over it and the products of combustion are absorbed^by 
 the water in the test-tube. Rinse out the combustion tube and 
 the delivery tube into the solution in the test-tube, filter and test 
 the filtrate with litmus. Make the following tests on this 
 solution : 
 
 71 
 
72 PHOSPHORUS, ARSENIC, ANTIMONY, ETC. 
 
 b. To a small portion add a few drops of potassium perman- 
 ganate solution and warm. 
 
 c. To another portion add silver nitrate solution, then sodium 
 bicarbonate (NaHC0 3 ) till a precipitate becomes permanent, 
 then warm. 
 
 d. Add a small portion to albumin solution acidified with 
 acetic acid. What do each of these tests indicate? What 
 materials are thus shown to be present in the solution? 
 
 e. Evaporate the rest of this solution in a porcelain dish, adding 
 nitric acid till brown fumes cease to be given off. (What are they ? 
 Explain their formation.) Evaporate till acid fumes are no 
 longer given off, then heat with the direct flame until there is 
 no further change, cool, and dissolve in water. Test this solu- 
 tion with potassium permanganate and silver nitrate solutions 
 as above, and explain the results. What is now present in the 
 solution? Allow the solution to stand for several days and test 
 again with silver nitrate. Explain. 
 
 56- The Preparation of Primary, Secondary and Tertiary 
 Orthophosphates, and their Decomposition Products on Heat- 
 ing. Quantitative. 
 
 a. Arrange twelve test-tubes in groups of three, and to the 
 first of each group add a solution of primary sodium phosphate 
 (NaH 2 P0 4 ); to the second, secondary sodium phosphate 
 (Na 2 HP0 4 ); to the third, tertiary sodium phosphate (Na 3 P0 4 ), 
 and label each tube. To each of the tubes in the first group add 
 a few drops of alizarine green, to the second group add phenol- 
 phthalein, to the third congo red, and to the fourth methyl orange. 
 Record the colors formed in tabular form, noting at the same 
 time the colors of each indicator in a solution of a strong acid 
 and of a strong alkali. Which indicator should be used in the 
 preparation of each of the sodium salts of orthophosphoric 
 acid by neutralization of the acid with sodium hydroxide? 
 Get the approval of an instructor on this point before 
 proceeding. 
 
 b. Prepare 100 c.c.of an approximately formular solution of 
 phosphoric acid by diluting the concentrated acid (see page 7 
 
ARSENIC 73 
 
 for the table of per cent composition and density) and an equal 
 amount of formular sodium hydroxide solution from the solid 
 sticks. 
 
 c. Into each of three evaporating dishes measure accurately 
 (pipette) 10 c.c. of the acid solution after it has been well mixed, 
 then neutralize the acid in each dish with the sodium hydroxide 
 solution measured from a burette, and using the selected in- 
 dicators, so that the primary, secondary and tertiary salts will 
 be formed. Record the volumes of alkali required in each case. 
 What ratio do these amounts bear to each other? What 
 ratio should they bear, theoretically? Evaporate the solu- 
 tions on the water bath, then heat the dry salts with the 
 direct flame until all moisture is removed. What salt is 
 now present in each case? Dissolve each of these three 
 salts in water and test each solution with silver nitrate solution, 
 and with albumin solution acidified with acetic acid. Do these 
 tests confirm your conclusions regarding the salts present? 
 
 d. Devise and perform tests to show that most phosphates are 
 insoluble in water but soluble in acids. 
 
 57. Arsenic. 
 
 a. Place about a gram of arsenopyrites in the center of a 
 hard glass tube such as was used in preparing the oxides of 
 phosphorus (Exp. 55), having one end closed with a cork, and 
 the other carrying a small bore delivery tube. Heat the pyrite, 
 when a sublimate will collect on either side of the hot part of 
 the tube. What is it? After a considerable sublimate has 
 collected allow the tube to cool, remove the residue of pyrite 
 and connect the combustion tube with a test-tube or flask, and 
 suction as in Exp. 55, Fig. 17, and burn one of the two deposits 
 (save the other for c) collecting the products in the test-tube 
 as before. Compare the solubility of arsenic oxides with that 
 of phosphorus oxides. Filter the solution of the oxides (keep 
 the insoluble residue for b) and test a small portion of the fil- 
 trate for sulfur dioxide by adding hydrochloric acid then barium 
 chloride a precipitate here would indicate what substance ? 
 
74 PHOSPHORUS, ARSENIC, ANTIMONY, ETC. 
 
 If any precipitate appears filter, then add potassium perman- 
 ganate solution to the filtrate and heat. A precipitate here is 
 due to what substance? Explain its presence. 
 
 b. Wash the residue of the arsenic oxide several times with 
 distilled water, then punch a hole in the bottom of the filter 
 paper and wash the oxide into a beaker with 20-30 c.c. of dis- 
 tilled water. Heat this to boiling to dissolve as much of the 
 oxide as possible, filter if the solution is not clear and test the 
 solution for sulfur dioxide as before. Is a precipitate of barium 
 sulfate formed? Is the potassium permanganate decolorized? 
 Explain. To a small portion of the solution add silver nitrate 
 then sodium carbonate solution till a permanent precipitate 
 forms (one drop is usually enough) and boil. Is the precipi- 
 tate silver carbonate? Devise and perform an experiment to 
 answer this question. What is the precipitate and what is its 
 reaction on boiling? Compare with the similar compound of 
 phosphorus. 
 
 Into the remainder of the solution pass hydrogen sulfide from 
 a clean delivery tube (if the containers and solution have not 
 been kept perfectly clean and free from electrolytes this will 
 be a failure) and note the change in color. This should give 
 a colloidal solution of arsenic sulfide. Compare the physical 
 properties of this solution with those of a true solution. Divide 
 the solution into two portions to one add hydrochloric acid, 
 to the other sodium chloride solution and explain the change. 
 Would it be classified as a chemical or physical change? 
 
 c. Dissolve the other portion of the arsenic sublimate (a) in 
 concentrated nitric acid and evaporate till the excess of acid is 
 removed (hood), add water and to a small portion of the solu- 
 tion, filtered if not clear, add silver nitrate and sodium car- 
 bonate solutions till a permanent precipitate forms. Why is 
 the sodium carbonate necessary? What is the precipitate? 
 Boil. Compare with the silver arsenite precipitate prepared in 
 6. Evaporate the rest of the solution to dryness and heat till 
 all moisture is removed. What remains? Compare with the 
 effect of heating phosphoric acid. 
 
BISMUTH 75 
 
 58. Antimony. 
 
 Treat 1-2 grams of stibnite (what is it ?) with concentrated 
 hydrochloric acid, warming in a covered evaporating dish until 
 the reaction becomes slow. Note the odor of the gas evolved. 
 What is it? Remove the solution from the excess of stibnite 
 and add an equal volume of water. If a white precipitate 
 forms (what is it?) add hydrochloric acid till it dissolves. If 
 the solution turns yellow or a yellow precipitate appears (due 
 to what?) boil till the color disappears. Make this solution 
 alkaline with ammonium hydroxide, filter out the precipitate 
 and wash with distilled water. What is the precipitate? Boil 
 a small portion of the precipitate with distilled water and test 
 with litmus. Is the hydroxide soluble? To two other por- 
 tions of the precipitate add sodium hydroxide solution and 
 hydrochloric acid respectively. Draw conclusions as to the 
 acidic or basic character of the hydroxide. Compare with the 
 character of the oxides of arsenic and phosphorus. 
 
 59. Bismuth. 
 
 Add 2-3 c.c. concentrated nitric acid to about 0.5 gram of 
 metallic bismuth, boil the solution till nitric oxide fumes are 
 no longer evolved and allow to crystallize by cooling and spon- 
 taneous evaporation. Heat a small portion of the crystals in 
 a dry test-tube, noting the gas evolved. Add a second small 
 portion to water and test with litmus. Account for the acid 
 solution and for the insoluble substance. Add nitric acid to 
 this until solution is complete, then prepare bismuth hydroxide 
 by adding ammonium hydroxide to the solution. Filter and 
 wash once with distilled water. Treat part of the precipitate 
 with a 25 per cent solution of sodium hydroxide, and part with 
 dilute hydrochloric acid. Is bismuth hydroxide basic or acidic? 
 Indicate its ionization products and the effect of the addition 
 of sodium hydroxide or hydrochloric acid on the ionization. 
 
CHAPTER XI 
 
 CARBON 
 
 60. Carbon. 
 
 a. Fill a crucible about one-fourth full of cane sugar, and heat 
 until the mass solidifies while hot. testing from time to time to 
 determine whether the gas evolved burns. Heat 4-5 grams of 
 copper oxide in a crucible to remove water, and when cool mix 
 with some of the sugar charcoal in a dry, hard glass test-tube 
 and heat, while watching for the condensation of moisture in 
 the cool part of the tube. Pass some of the gases evolved into 
 lime water, and explain the result. To what element in the 
 charcoal is the moisture due? Is charcoal often pure carbon? 
 Name the forms in which pure carbon occurs. Burn the rest 
 of the charcoal, observing the rate and ease of combustion. 
 Is carbon an active element? Does it burn with a flame? 
 
 b. Heat gently about 2 grams of soft coai in a crucible covered 
 with an inverted funnel, noting the formation of a gas (does it 
 burn?) and of tar. When combustible gases are no longer 
 given off cool and observe the appearance of the coke formed. 
 Is it pure carbon? To answer this apply the test used in a 
 on one portion of the coke and place the crucible containing the 
 rest of the coke in an inclined position and heat until the carbon 
 has all burned away. What is left? What are some of the 
 points to be considered in buying coal for heating purposes? 
 
 61. Carbon Dioxide. 
 
 Add dilute hydrochloric acid to marble in a generating flask 
 (Fig. 7) and pass the gases evolved into lime water. What is 
 formed? Attach a wash bottle to the generator so that the gas 
 bubbles through water (to remove hydrochloric acid which 
 might be carried over mechanically), then pass the gas into a 
 
 76 
 
CARBON MONOXIDE 77 
 
 flask containing about 100 c.c. of water until the solution is satu- 
 rated. Test the solution with litmus paper, taste it. What is 
 present ? Keep this solution (stoppered) for Exp. 64. Determine 
 whether the gas is lighter or heavier than air, and fill a bottle of 
 the gas by air displacement. Insert into this a burning taper. 
 Pour the gas from one bottle into another and demonstrate that 
 it has really passed from one to the other. Direct the stream 
 of gas from the generator upon the flame from a burning match 
 or taper, and explain. What commercial application is made 
 of this property? 
 
 62. Determination of the Density of Carbon Dioxide. Quan- 
 titative. 
 
 Which of the methods previously used for density determina- 
 tions can be applied to carbon dioxide? Arrange the apparatus 
 and proceed with the determination after getting the approval 
 of an instructor. 
 
 63. Carbon Monoxide. Poison. Do not Let the Gas Escape 
 into the Room. 
 
 a. Mix about 3 grams of solid oxalic acid with an equal 
 volume of concentrated sulfuric acid in a generating flask and 
 heat, collecting three bottles of the gas by displacement of water. 
 To stop the evolution of gas allow the generating flask to cool. 
 
 b. Add lime water to the first bottle and shake. What is 
 formed? 
 
 c. Invert the second bottle as quickly as possible over a beaker 
 of dilute potassium hydroxide solution and allow to stand until 
 absorption is complete, then cover, set the bottle upright 
 without pouring out the water that was absorbed and apply a 
 light. What is this gas? What were the relative proportions 
 of carbon dioxide and carbon monoxide in the gas mixture? 
 
 d. Apply a light to the third bottle of the gas mixture, and 
 explain the result. 
 
 64. Hard Water. 
 
 a. Dilute some lime water with an equal volume of distilled 
 water and pass in carbon dioxide until the precipitate first 
 
78 CARBON 
 
 formed (what is it?) redissolves (what is formed?). Filter 
 the solution if it is not entirely clear, and divide into five por- 
 tions. Boil one portion. To the other four add sodium car- 
 bonate, lime water, soap and alum solutions respectively. Draw 
 conclusions as to the efficiency of these reagents and of boiling 
 as softening agents? 
 
 b. Pour 20 c.c. of the carbonic acid solution prepared in 
 Exp. 61 over a small amount of powdered marble or chalk, and 
 allow to stand over night, then filter and test the filtrate by 
 boiling, with sodium carbonate, etc. What was present in the 
 solution? What is the importance of this reaction in nature? 
 
 65. Soap. 
 
 a. Dissolve about 0.5 gram of soap by heating in 25 c.c. of 
 distilled water and test the solution with litmus and explain. 
 Add solutions of sodium and calcium salts to different portions 
 of this soap solution. Which would be more objectionable in 
 water for laundry purposes, sodium or calcium bicarbonate? 
 
 b. Place equal quantities of animal charcoal in each of two 
 test-tubes; to one add 10 c.c. of hot soap solution, to the other 
 10 c.c. of hot water, filter each and compare the filtrates. The 
 soap solution tends to combine with the carbon to form an 
 adsorption compound which reduces the tendency for the carbon 
 to stick to other materials, here the filter paper. This property 
 partly explains the action of soap in cleansing. 
 
 c. To 10 c.c. of soap solution add dilute hydrochloric acid to 
 acid reaction. Warm gently until the white precipitate melts 
 and floats on the surface. What is it? 
 
 66. Methane. 
 
 a. Grind together one part of sodium acetate with four parts 
 of soda lime (quick lime and sodium hydroxide) and place the 
 mixture in a hard glass test-tube fitted with a delivery tube. 
 Incline the tube so that water driven from the contents on heat- 
 ing will not run back upon the hot part of the glass, and warm, 
 collecting the gas evolved by displacement of water. 
 
HYDROCARBONS 79 
 
 b. Shake one bottle of the gas with a few cubic centi- 
 meters of potassium permanganate solution. 
 
 c. Shake another with bromine water. 
 
 d. Burn a third and test the products of combustion remain- 
 ing in the bottle for carbon dioxide. 
 
 e. Light a few bubbles of the gas as they escape from the 
 water and note the nature of the flame. 
 
 67. Ethylene. 
 
 Add 9 c.c. alcohol slowly to 60 c.c. concentrated sulfuric acid, 
 stirring constantly. Pour this mixture into a flask fitted with 
 a delivery tube leading under water, and clamped so as to rest 
 on a sand bath. Add a little sand to reduce foaming and heat 
 gentry until there is a steady evolution of gas (avoid over- 
 heating as the mixture foams considerably) and collect bottles 
 of the gas, repeating all the tests applied to methane and 
 comparing the results. 
 
 68. Acetylene. 
 
 a. Place 15 grams of calcium carbide in a flask fitted with a 
 separatory funnel and delivery tube which is connected with a 
 hard glass combustion tube open at each end and fitted with a 
 delivery tube leading under water, Fig. 18. Generate the gas 
 by allowing water from the separatory funnel to drop upon the 
 calcium carbide, and collect bottles of the gas to which all the 
 tests under methane should be applied. Compare with methane 
 and ethylene. 
 
 6. After all the air has been driven out from the generating 
 flask and delivery tubes heat the hard glass tube while still 
 generating acetylene, and observe the deposit on the glass, and 
 the changed appearance of the gas evolved, and explain. 
 
 69. The Oxidation Products of Alcohol. 
 
 Add 2 c.c. of concentrated sulfuric acid to a mixture of 8 c.c. 
 water and 2 c.c. 50 per cent alcohol, mix thoroughly then add 
 a small crystal of potassium dichromate and warm. Test the 
 
80 
 
 CARBON 
 
 gas evolved by holding in it strips of moist litmus paper ; and note 
 its odor. This is acetaldehyde. 
 
 To 1 to 2 c.c. of concentrated sulfuric acid add about 5 drops of 
 50 per cent alcohol, mix, then add a small crystal of potassium di- 
 
 ll 
 
 FIG. 18. 
 
 chromate and warm till the reaction begins (care !) . Test the gas 
 evolved with moist litmus paper and note its odor. What is it? 
 Write equations for both reactions and explain the change in 
 color of the solutions. 
 
CHAPTER XII 
 
 SILICON 
 
 70. Silicon Tetrafluoride. 
 
 Grind together 1 gram of silicon dioxide (sand) and 3 grams 
 of calcium fluoride, moisten with concentrated sulfuric acid and 
 heat in a test-tube fitted with a delivery tube leading just above 
 the surface of water. What is the precipitate formed on the 
 surface of the water? Explain the fumes above the water. 
 Test the water solution by filtering, then adding to the filtrate 
 dilute hydrochloric acid and barium chloride solution, and warm- 
 ing. What is formed? Write all reactions. 
 
 71. Silicic Acid. 
 
 a. Grind together in an iron crucible 1 gram of sand and 3 
 grams of sodium hydroxide (solid) and heat until the mass 
 which at first fuses, turns solid. Cool, dissolve in 25 c.c. water 
 and filter. To one portion of the filtrate (what is it?) add con- 
 centrated or dilute hydrochloric acid very slowly. Pour another 
 portion, about 6 c.c., into a test-tube containing about 6 c.c. of 
 concentrated hydrochloric acid. Has there been any difference 
 in the chemical reaction in these two tubes? Explain the 
 difference in appearance. Evaporate part of the clear solution 
 of silicic acid in a porcelain dish and heat the dry residue with 
 the direct flame, then cool and determine whether it is still com- 
 pletely soluble in water. 
 
 b. Dialyze the rest of the silicic acid solution as follows: 
 Prepare a collodion sack for dialysis by wetting the inside walls 
 of a 50 c.c. beaker which has been carefully cleaned and dried, 
 with a solution of collodion in ether and allowing the ether to 
 evaporate, then when the odor of ether has disappeared remove 
 the sack from the beaker (loosen the edge of the collodion from 
 
 81 
 
82 
 
 SILICON 
 
 the glass with a knife, then the sack may be withdrawn with the 
 finger, or by allowing water to run in between the glass and the 
 collodion) and fill with the solution of colloidal silicic acid. 
 Suspend this sack in a beaker of distilled water so that the level 
 of the water is slightly lower than the level of the solution in the 
 sack (Fig. 19) and change the dialysate (the solution in the 
 
 FIG. 19. 
 
 beaker, consisting of materials which have dialyzed through 
 the membrane) frequently until it gives only a very faint test 
 for chlorides after standing in contact with the sack for about 
 an hour. Test portions of the dialyzed solution (the solution 
 remaining in the sack) with barium chloride and copper sulfate 
 solutions. Would barium or copper silicates be soluble if 
 formed ? Is the silicic acid ionised in this solution ? Evaporate 
 another portion to dryness, heat with the direct flame and 
 determine whether the residue is soluble in water. Has it 
 undergone a chemical change? 
 
CHAPTER XIII 
 BORON 
 
 72. Preparation of Boric Acid from Borax. 
 
 Saturate 50 c.c. of distilled water with borax, test the solution 
 with litmus and account for the result. Add dilute sulfuric acid 
 to this solution until no further precipitate forms, then cool and 
 filter the cold solution. Wash once or twice with cold distilled 
 water then purify the precipitate by redissolving in the least 
 possible amount of boiling water and cool under running water. 
 Filter and wash the precipitate with cold water till the wash 
 water is free from sulfates, then dry thoroughly by pressing 
 between folds of filter paper and then warming on the radiator 
 or in a drying oven at a temperature below 100 C. What is the 
 residue? Dissolve some in water and test the solution with 
 litmus. Mix a small portion with an equal quantity of sodium 
 chloride and heat the mixture in a test-tube. What is formed? 
 Heat a small portion in a dry test-tube. Do these tests prove 
 whether the material is an acid or an acid anhydride? 
 
 73. Boric Acid Anhydride or Boron Trioxide from the Acid. 
 Quantitative. 
 
 Carefully clean and weigh a porcelain crucible and place in it 
 about 1 gram of the well dried acid from the previous experi- 
 ment (Exp. 72) and weigh accurately. Cover the crucible and 
 heat, at first gently and then with the full heat of a bunsen burner. 
 Cool and weigh, and repeat until the crucible and contents come 
 to constant weight. Calculate the per cent of water in the acid, 
 and the number of molecules of water to one of the acid 
 anhydride. 
 
 74. The Preparation of Metallic Boron. 
 
 Heat some of the prepared boric acid (Exp. 72) in a porcelain 
 crucible with the direct flame till all water is removed, then mix 
 
 83 
 
84 BORON 
 
 about 1 gram of the anhydride with 1.5 grams of magnesium 
 powder. Place a small portion of this mixture in a crucible 
 and heat till glowing ceases, then add gradually the rest of the 
 mixture, heating continuously. Grind the product in a mortar, 
 wash it with dilute hydrochloric acid then with hot concentrated 
 hydrochloric acid diluted 1 :1 with water, and finally with dis- 
 tilled water until the filtrate comes through brown, due to the 
 formation of a colloidal solution of boron, and indicating that 
 electrolytes are not present in appreciable amounts. Test the 
 solubility of this boron in concentrated nitric, hydrochloric and 
 sulfuric acids. Fuse a small portion with an equal quantity of 
 solid potassium hydroxide. What is formed? Dissolve the 
 fused mass in as little water as possible and acidify with dilute 
 sulfuric acid. What is formed? How may tertiary sodium 
 borate (Na 3 B0 3 ) be prepared? Is it stable in solution? 
 
 75. Berates of the Heavy Metals. 
 
 Will borax or boric acid precipitate copper borate from copper 
 sulfate solution? Is copper borate soluble in water? To 
 answer this question prepare a borax bead by fusing borax on 
 the loop of a platinum wire, dip the bead into copper oxide and 
 heat until a transparent colored bead is formed. If the bead 
 does not become clear after heating several moments it is be- 
 cause too much copper oxide has been added. What copper 
 compound is present in the bead? Break up the bead and test 
 its solubility in water. Boil the water supension and note the 
 change in color of the precipitate, which is due to the formation 
 of copper oxide from the hydroxide first formed. Explain 
 the presence of copper hydroxide on treating the bead with 
 water, on the basis of the ionization theory. 
 
CHAPTER XIV 
 
 METALS AND NON-METALS 
 
 76. Hydrolysis. 
 
 Dissolve crystals of sodium chloride, potassium chloride, 
 ammonium chloride, cupric chloride, ferric chloride, stannous 
 chloride and phosphorus pentachloride each in a separate 
 beaker in about 5 to 10 c.c. of water. Test the solutions with 
 litmus, then boil each, covering the beaker with a watch glass 
 on which is placed moist red and blue litmus paper so that it 
 will come in contact with the vapor from the boiling solution. 
 Note and account for the effect of the vapor on the litmus in 
 each case, then allow the solutions to evaporate to dryness on 
 the water bath, dissolve the dry residue in very dilute nitric 
 acid and test the solution for chlorides. Which of these salts 
 have undergone no hydrolysis? Which have been partially 
 hy droly zed ? Which have been completely hy droly zed ? Define 
 hydrolysis. What is the effect of water on the chlorides of 
 metals ? Of non-metals ? Do the results of this experiment lead 
 to the conclusion that there is a sharp distinction between 
 metals and non-metals? 
 
 77. lonization. 
 
 a. Make up 100 c.c. each of thrice normal solutions of sulfuric, 
 hydrochloric, acetic, phosphoric and oxalic acids. To deter- 
 mine the relative strength, or degree of ionization of these acids 
 place 5 c.c. portions of each acid in separate test-tubes, 
 label each tube and place them all in a beaker containing 
 water, so that they will not suffer any great change of tempera- 
 ture. Now into each tube drop a small granule of zinc and note 
 the rate of the reaction, j udging by the rate of evolution of hy- 
 drogen in each tube. Name the acids in the order of their 
 
 85 
 
86 METALS AND NON-METALS 
 
 strength, then heat the water bath till it reaches about 50 and 
 again note the rate of the reaction, and record. If the weakest 
 reaction were considered the standard or unit of strength, 
 what values would you give the other acids? Compare your 
 values with those in the table of degrees of ionization. Repeat 
 this experiment using small chunks of marble instead of zinc, 
 and arrange the acids according to their strength, as before. 
 How does this order agree with that arranged with reference to 
 their action on zinc? 
 
 b. Make a saturated solution of potassium chromate. and 
 place equal quantities in each of six test-tubes. Now add to one 
 of the test-tubes, drop by drop, the acid found to be weakest in 
 the previous test, until a just noticeable change in color has 
 occurred, then add to each of the other test-tubes an equal num- 
 ber of drops of each of the acids, label each, and to the sixth 
 tube add an equal number of drops of concentrated sulfuric acid 
 (this represents the greatest possible change) . Shake each tube 
 and arrange in the order of their depth of color, and compare 
 with the order of the acids with respect to their activity on zinc 
 or marble. 
 
 c. Repeat this procedure, using a solution of 6 grams po- 
 tassium iodide and 1 gram potassium bromate per liter instead 
 of the solution of potassium chromate. In this case each of the 
 tests will finally come to the same color, but the weakest acid 
 will bring about the change most slowly, so that they may be 
 arranged according to either of two variables, i.e., the depth of 
 color at a fixed time after adding the acid, or the length of 
 time required for each tube to come to a definite depth of color. 
 Record the acids in the order of one of these variables, and com- 
 pare with the order found in a and b. 
 
 78. Indicators. 
 
 a. Dilute 1 c.c. of concentrated hydrochloric acid with 9 c.c. 
 of water, mix thoroughly and add one drop of this diluted acid 
 to 100 c.c. of distilled water (A). Mix thoroughly and add 
 equal portions to four test-tubes, then add to each test-tube a 
 
INDICATORS 87 
 
 few drops of one of the indicators, methyl orange, methyl red, 
 congo red and litmus, the acid and alkali colors of each of which 
 should be previously determined. Dilute 50 c.c. of this acid 
 (A) with 50 c.c. of distilled water, mix thoroughly and test 
 again with each of the four indicators, and repeat this diluting 
 and testing until each of the four indicators have changed color. 
 Which indicator changes color in the greatest hydrogen ion 
 concentration? which in the least? 
 
 b. Make up a solution of ammonium hydroxide in exactly the 
 same way that the hydrochloric acid solution was prepared, 
 using the concentrated ammonium hydroxide, sp. gr. 0.90, 
 and test the successive dilutions of this with phenol phthalein, 
 tropaolin O and litmus. Which indicator changes color in 
 the greatest hydroxyl ion concentration? which in the least? 
 
 c. To two portions of 10 c.c. of normal sodium acetate add 
 methyl orange and phenolphthalein respectively as indicators 
 (is the reaction of sodium acetate acid or alkaline?) then titrate 
 each with normal acetic acid and record the quantities of acid 
 used to neutralize the solution, with each indicator. Explain 
 why these quantities shoud be different. Now repeat the pro- 
 cedure titrating with normal hydrochloric acid instead of acetic, 
 record the quantities used and explain the different effects of 
 the two acids. In neutralizing sodium hydroxide with acetic 
 acid which of these two indicators should be used? In neutral- 
 izing with hydrochloric acid would it make much difference 
 which was used? 
 
 d. Test a solution of ammonium sulfate with methyl orange 
 and phenolphthalein, and titrate equal volumes of the solution 
 with normal ammonium hydroxide or normal sulfuric acid, 
 according as the original solution reacts acid or alkaline, using 
 these two indicators. What indicator should be used in neu- 
 tralizing a strong acid, as sulfuric acid, with a weak base, as 
 ammonium hydroxide? Titrate equal quantities of normal 
 ammonium hydroxide with normal acetic acid, using methyl 
 orange and phenolphthalein as indicators. Will these indi- 
 cators serve to show the neutral point in this case? Why? 
 
88 METALS AND NON-METALS 
 
 79. Indicators. Quantitative. 
 
 a. By titrating a mixture of sodium hydroxide and sodium 
 bicarbonate with normal hydrochloric acid it is possible to 
 titrate only the sodium hydroxide, if the indicator is chosen so 
 that as soon as carbonic acid is formed it shows acid reaction, 
 and with another indicator which is not affected by so low a 
 hydrogen ion concentration as is supplied by carbonic acid, 
 but is affected as soon as free hydrochloric acid exists in the 
 solution, it is possible to determine the total sodium hydroxide 
 plus sodium bicarbonate in the solution. Which indicator 
 should be used for each titration? Test your decision by try- 
 ing a solution of sodium bicarbonate alone with the first in- 
 dicator, and determine how much normal hydrochloric acid is 
 required to give acid reaction, and by trying the second indi- 
 cator to determine that carbonic acid does not give acid reac- 
 tion, but that free hydrochloric acid does. 
 
 6. Titrate equal volumes (pipette) of a solution of these two 
 salts with normal hydrochloric acid, using the two indicators 
 chosen, and from these two titrations calculate the amounts of 
 sodium hydroxide and sodium bicarbonate present in the 
 mixture. In the original solution could sodium hydroxide and 
 sodium bicarbonate exist together? What would be formed? 
 Is there an indicator which would give an acid reaction with 
 sodium bicarbonate, so that when a mixture of sodium hydrox- 
 ide and sodium carbonate were titrated only the sodium hydrox- 
 ide would be neutralized? Write all the reactions which occur 
 when sodium hydroxide and sodium bicarbonate are mixed 
 in solution, and when the mixture is titrated, first with one indi- 
 cator and then with the other. From the results of the titra- 
 tion above calculate the amount of sodium carbonate in the 
 given mixture. 
 
CHAPTER XV 
 THE ALKALI METALS 
 
 80. The Transition of Sodium Sulfate from the Anhydrous 
 to the Hydrated Salt. 
 
 Add 30 grams of anhydrous sodium sulfate (if only 
 the hydrated salt is available calculate the quantity re- 
 quired to give 30 grams of the anhydrous salt) to 40 c.c. of 
 water heated to 50 C., and hold the mixture at that tempera- 
 ture for about five minutes. Is the solid phase the hydrated 
 or the anhydrous salt? Can you suggest an experimental proof 
 for your answer to this question? Now cool by immersing the 
 beaker in water at 15 to 20 C. and record the temperature of the 
 solution every half minute for ten minutes, stirring it constantly. 
 If the temperature falls below 25 drop in a few crystals of the 
 hydrated salt. Make a plot to represent the change in tempera- 
 ture with the time, letting the horizontal axis represent tem- 
 perature and the vertical axis time, and explain the sudden 
 break in the curve. Is the change from the hydrated to the 
 anhydrous salt endothermic or exothermic? 
 
 81. Sodium Carbonate by the Solvay or Ammonia Process. 
 
 (Adapted from Blanchard's Synthetic Inorganic Chemistry.) 
 a. To 50 c.c. of concentrated ammonium hydroxide (0.90 
 specific gravity) add 150 c.c. of distilled water. Place in a 
 flask and add 60 grams of sodium chloride free from lumps. 
 Shake until the salt is nearly or quite dissolved, and filter the 
 solution, if it is not perfectly clear. Pass a delivery tube through 
 one hole of a two-hole, tightly fitting stopper placed in a 300 
 c.c. flask. Provide a plug for the other hole. Let the tube 
 dip into the solution which is placed in the flask, and pass in 
 carbon dioxide gas from a Kipp generator until all the air has 
 
 89 
 
90 THE ALKALI METALS 
 
 been displaced from the flask; then close the flask and allow the 
 gas to pass in as fast as it will be absorbed. What is 
 formed? Occasionally as the action seems to slacken, loosen 
 the plug for a moment. Shake the flask frequently. It will 
 take several hours for the solution to absorb sufficient carbon 
 dioxide, and it may be left over night connected with the 
 generator. When no more gas can be absorbed pour the 
 mixture from the flask onto a fluted filter paper, wash 
 the precipitate with a small amount of cold water, and 
 allow it to dry in the air. Look up the solubilities of the 
 substances started with and the final products of this 
 reaction and explain upon what the success of the process 
 depends. Why could not ammonium chloride be used instead 
 of ammonium hydroxide? What remains in the filtrate from 
 the bicarbonate? Would the process be financially a success 
 if this filtrate was discarded ? How may the ammonium hydrox- 
 ide be regenerated from it and what is the final waste product ? 
 
 b. Add hydrochloric acid to some of the solid bicarbonate 
 prepared above. Heat a small portion and test the gas evolved 
 for carbon dioxide. To the residue after heating add hydro- 
 chloric acid. What is the reaction on heating? 
 
 c. Heat about 1 gram of the bicarbonate till decomposition 
 is complete, dissolve the product in very little water, test with 
 litmus and compare with the effect of sodium bicarbonate 
 solution on litmus. Add a solution of lime water (calcium 
 hydroxide) till there is no further precipitation, and filter the 
 precipitate? What is in solution? 
 
 82. Potassium Nitrate from Sodium Nitrate and Potassium 
 Chloride. (Adapted from Blanchard's Synthetic Inorganic 
 Chemistry.) 
 
 Dissolve 50 grams of sodium nitrate and 44 grams of potas- 
 sium chloride in 100 c.c. of water and evaporate on the water 
 bath in a porcelain dish, adding the solution in portions if the 
 dish will not hold it all at once, till the solution is half its 
 original volume. Without letting the liquor cool separate it 
 
POTASSIUM NITRATE 91 
 
 from the crystals which have formed during the evaporation. 
 This may be done by decantation or by filtering through a fluted 
 filter paper which has been wet with hot water. Allow the 
 filtrate to cool by holding the beaker containing it under run- 
 ning water, and filter out the second crop of crystals. What 
 four salts may be present in the solution after potassium chlo- 
 ride and sodium nitrate have been mixed together? Look up 
 the solubilities of each of these salts and determine which would 
 separate out first from a solution containing equi-molecular 
 amounts of each, on evaporating at 100. Test the first crop 
 of crystals to determine if your reasoning is correct. If half 
 the sodium and chloride ions in the original mixture were 
 removed, what would be left in the solution and in what pro- 
 portions? Considering their solubilities at high and low tem- 
 perature, and the relative amounts present, which salt would 
 you expect would separate out on cooling the solution? Devise 
 and apply a test on the second crop of crystals to determine 
 whether your conclusion is correct. Why is potassium nitrate 
 used for the manufacture of gunpowder rather than sodium 
 nitrate? 
 
CHAPTER XVI 
 COPPER AND SILVER 
 
 83. Copper and Cupric Compounds. 
 
 a. Treat bright copper filings with dilute sulfuric and dilute 
 hydrochloric acids. Does the metal dissolve in either case? 
 Pour off the dilute acids and replace them with concentrated 
 acids, warm, and note the effect in each case. Test the effect 
 of both dilute and concentrated nitric acid on copper and ac- 
 count for the effect of each acid on the copper. Will copper 
 displace hydrogen from an acid? 
 
 6. Heat some copper filings in an open crucible and note the 
 change in appearance. Test the product for solubility in dilute 
 sulfuric and dilute hydrochloric acids, and account for the dif- 
 ference in behavior of copper before and after heating in air. 
 
 c. Add ammonium hydroxide drop by drop to about 5 c.c. of 
 copper sulfate solution, shaking after each addition, until the 
 mixture smells strongly of ammonia. Account for every change 
 which occurs. To another portion of copper sulfate solution 
 add sodium hydroxide solution till the mixture is alkaline, then 
 heat, and account for every change which occurs. 
 
 84. Cuprous Compounds. 
 
 a. Prepare a solution of 3.5 grams of copper sulfate in 50 c.c. 
 of water, and another solution of 17 grams of Rochelle salt 
 (sodium potassium tartrate) and 7 grams of potassium hydrox- 
 ide (or its equivalent of sodium hydroxide) in 50 c.c. of water. 
 Mix these solutions and heat to boiling then add a solution of 
 glucose until all the copper is precipitated. How can that 
 point be recognized? Filter and wash the precipitate (what is 
 it?) with warm water several times, then allow to dry. 
 
 6. To a small portion of the precipitate add dilute hydro- 
 chloric acid. What is formed? Test different portions of this 
 
 92 
 
CUPROUS COMPOUNDS 93 
 
 white precipitate for solubility in concentrated hydrochloric 
 acid and in ammonium hydroxide. Is the cupric ion present 
 in either of these solutions? State the reason for your answer. 
 Pour the concentrated hydrochloric acid solution into an excess 
 of water and explain. 
 
 c. Add ammonium hydroxide to a small portion of the red 
 cuprous oxide and explain the color changes. 
 
 d. Add dilute sulfuric acid to another portion of the cuprous 
 oxide. Note the color of the solution and of the insoluble 
 residue. Devise and apply tests to prove that this residue is not 
 cuprous oxide, but is metallic copper. Is the cuprous or cupric 
 ion present in the solution? What is the oxidizing and what the 
 reducing agent in this reaction? 
 
 e. To a few cubic centimeters of copper sulfate solution add 
 an excess of potassium iodide solution, filter and wash the 
 precipitate first with a little potassium iodide solution (why?), 
 then with water. Explain the color of the filtrate. To one 
 portion of the precipitate add dilute potassium hydroxide and 
 warm. Is this the oxide of cupric or cuprous copper? To 
 another portion of the precipitate add ammonium hydroxide. 
 What is formed? 
 
 85. Electromotive Series. 
 
 Place the following mixtures in test-tubes, taking especial 
 pains to see that the metal has a bright, shiny surface, i.e., is 
 not covered with an oxide. Observe the changes occurring, 
 and complete the chemical equations 
 
 Cu+ZnS0 4 solution. Pb-f-ZnS0 4 solution. 
 
 Cu + Pb(N0 3 ) 2 solution. Pb + Cu(N0 3 ) 2 solution. 
 
 Cu+Eg(N0 8 )j solution. Pb + Hg(N0 3 ) 2 solution. 
 
 Hg+ZnS0 4 solution. 
 
 Hg + Pb(N0 3 ) 2 solution. 
 
 Hg + CuS0 4 solution. 
 
 Zn + Pb(N0 3 ) 2 solution. Pb + HCl (dilute). 
 
 Zn+Cu(N0 8 ) 2 solution. 
 
 Zn + Hg(N0 3 ) 2 solution. 
 
94 COPPER AND SILVER 
 
 These reactions are not instantaneous and should be allowed to 
 continue during the time of at least one laboratory period. 
 Arrange the four elements in the order of their " solution 
 tension" or tendency to change from the free to the ionic con- 
 dition. Refer to the last part of Exp. 13 and find the place 
 for hydrogen in the series. 
 
 86. Silver Oxide. 
 
 a. Prepare silver oxide by adding an excess of potassium hy- 
 droxide to about 20 c.c. of silver nitrate solution, filter out the 
 oxide and wash, then test small portions of the precipitate for 
 solubility in ammonium hydroxide and dilute sulfuric acid. 
 Heat the remainder of the precipitate in a porcelain crucible, 
 and note the change in appearance and test small portions 
 for solubility in ammonium hydroxide and dilute sulfuric acid. 
 Explain the difference in conduct before and after heating. 
 Dissolve the remaining silver in concentrated nitric acid which has 
 been diluted with an equal volume of water, avoiding an excess 
 of the acid, then evaporate the solution till crystals begin to 
 form on the surface, then crystallize by cooling and spontaneous 
 evaporation. 
 
 b. Convert a small portion of the silver nitrate into silver 
 chloride, add ammonium hydroxide and explain. Devise and 
 apply a method for converting the rest to silver sulfate. Com- 
 pare the properties of the cuprous and silver ions. 
 
 87. Silver Chloride Emulsion. 
 
 Prepare a 1 per cent gelatin solution (gelatin will dissolve 
 readily on warming with water) and to a measured amount of 
 the solution (about 25 c.c.) add two drops of concentrated 
 hydrochloric acid and an excess of silver nitrate solution and 
 compare with the appearance of the silver chloride when the 
 same procedure is followed, using distilled water instead of the 
 gelatin solution. Allow the two preparations to stand for an 
 hour or more and again compare. What is the technical impor- 
 tance of silver chloride emulsions? 
 
CHAPTER XVII 
 CALCIUM, STRONTIUM AND BARIUM 
 
 88. Calcium Carbonate and Lime. 
 
 a. Heat about 2 grams of well ground marble in a hard glass 
 test -tube fitted with a stopper and delivery tube, and test the 
 gas evolved by passing it through lime water. Remove small 
 portions of the material in the test-tube at intervals of about fif- 
 teen minutes, and test with hydrochloric acid to determine the 
 rate of decomposition. How could marble be kept from decom- 
 posing at 800 ? After the heating has been continued for about 
 an hour test a portion of the residue in the test-tube for solubility 
 in water and for the reaction of the solution toward litmus, and 
 compare with the original marble. Account for the differences 
 noted. Filter out the insoluble' residue and pass carbon dioxide 
 into the filtrate. What is formed? Add a little quicklime to 
 water and note the temperature change, and the reaction of 
 the solution toward litmus. 
 
 89. The Relative Solubility of the Hydroxides of Calcium, Stron- 
 tium and Barium. Quantitative. 
 
 a. Make saturated solutions of barium, strontium and calcium 
 hydroxides by adding an excess of the solid to 50 c.c. of distilled 
 water and allowing to stand in stoppered flasks until the super- 
 natant liquid is clear. (Keep these solutions for Exp. 90.) 
 Now make 200 c.c. of an approximately 0.1 normal hydro- 
 chloric acid solution by diluting the concentrated acid (sp. gr. 
 1.16). Take 10 c.c. of each of the hydroxide solutions, measured 
 with a pipette, from the clear supernatant liquid. Neutralize 
 these with the acid from a burette, recording the exact quantity 
 of acid required in each case and using an indicator which will 
 not give the acid color in so slight a hydrogen ion concentration 
 
 95 
 
96 CALCIUM, STRONTIUM AND BARIUM 
 
 as comes from carbonic acid. Why? Calculate the relative 
 solubilities of these three hydroxides. If the acid is assumed to 
 be exactly 0.1 normal what are the absolute solubilities? 
 
 b. Titrate 100 c.c.of tap water in the same way. What is the 
 alkaline substance? Assuming that the alkali in the tap water 
 is lime, express the concentration of the tap water in terms of a 
 saturated solution of lime. (That is, if a saturated solution 
 was taken as unity, this would be some fraction of a saturated 
 solution.) 
 
 90. The Relative Solubilities of Calcium, Strontium and Barium 
 Bicarbonates. 
 
 Measure out 10 c.c. of each of the saturated solutions of cal- 
 cium, strontium and barium hydroxides prepared in Exp. 89 
 and pass carbon dioxide from a Kipp generator into each of these 
 solutions, adding water to each gradually, while the carbon 
 dioxide continues to pass through, until the precipitate first 
 formed just dissolves in the excess of water saturated with 
 carbon dioxide. (A small granular precipitate in the case of the 
 barium may be ignored.) Now measure the total volumes of 
 each of the solutions, and calculate the relative solubilities of 
 the bicarbonates from that of the hydroxides (Exp. 89) and 
 these measured volumes. 
 
 91. The Relative Solubilities of Calcium, Strontium and Barium 
 Sulfates. 
 
 Prepare calcium sulfate from a small piece of marble (can 
 this be done by adding sulfuric acid directly? Why?). Filter 
 and wash the calcium sulfate free from sulfuric acid (would you 
 test this with litmus or with barium chloride? Why?) then 
 make a saturated solution of this calcium sulfate and add the 
 clear solution to strontium chloride solution and warm. What is 
 precipitated? What remains in solution? Add the filtrate 
 from this to barium chloride solution and explain. Draw con- 
 clusions regarding the relative solubility of barium, strontium 
 and calcium sulfates. Explain why calcium sulfate precipi- 
 tates from solutions of calcium chloride and sulfuric acid more 
 
CALCIUM CHLORIDE 97 
 
 slowly than barium sulfate from solutions of barium chloride 
 and sulfuric acid. 
 
 92. Preparation and Properties of Calcium Chloride. 
 
 a. Prepare a solution of calcium chloride from about 25 
 grams of marble, taking care to avoid an excess of hydrochloric 
 acid, filter and crystallize by boiling until crystals begin to 
 separate, then allow to cool, remove the crystals from the 
 mother liquor, and dry them on filter paper. Boil down the 
 mother liquor and continue heating, with stirring, until the 
 solid residue has become quite white and opaque. What is the 
 difference between these two preparations? How much of the 
 crystallized salt would be equivalent to a gram of the 
 anhydrous ? 
 
 b. Add equivalent quantities of each of these salts (finely 
 ground) to 10 c.c. of water, taking the temperature before addi- 
 tion and after solution of the salt, and account for the difference. 
 
 c. Repeat b, using equal weights of finely chopped ice or 
 snow instead of water, and account for the difference in effect 
 of the two salts, and for the difference in effect as compared with 
 the previous test with water. Allow samples of each of the 
 preparations to stand in the open air and account for the change. 
 
CHAPTER XVIII 
 MAGNESIUM ZINC AND MERCURY 
 
 93. The Equivalent Weight of Magnesium. Quantitative. 
 
 Prepare an apparatus as shown in Fig. 20, in which an ordi- 
 nary test-tube is fitted with a two hole stopper, one hole of 
 
 FIG. 20 
 
 which carries a glass rod which has been bent over just a little 
 at its lower end, and after being put into the cork has been bent 
 
 98 
 
ZINC 99 
 
 at right angles at the other end, while the other hole carries a 
 tube bent at an acute angle, to connect with the gas burette. 
 Make the apparatus air tight, then fill the burette to the zero 
 mark with water, place about 10 c.c. of dilute hydrochloric 
 acid in the test-tube, and hang on the glass rod a weighed 
 quantity of magnesium ribbon (not more than 0.05 gram) or a 
 measured length, the weight of which can be calculated from a 
 factor, weight per centimeter, on the bottle. Connect the ap- 
 paratus, open the burette cock and read the volume (at atmos- 
 pheric pressure), drop the magnesium ribbon into the acid by 
 turning the glass rod, and after the magnesium is all dissolved 
 read the volume again, noting the temperature of the room and 
 the atmospheric pressure. Calculate the weight of hydrogen 
 evolved, and the equivalent weight of magnesium. 
 
 94. Magnesium Carbonate and Oxide. 
 
 Heat magnesium carbonate as calcium carbonate was heated 
 (Exp. 88), testing a small portion with hydrochloric acid every 
 ten minutes till decomposition is complete, add water and test 
 the suspension with litmus. Compare these with the same 
 experiments on calcium. Add a saturated solution of am- 
 monium chloride to the suspension of magnesium oxide, test 
 again with litmus and account for the changes. 
 
 95. The Properties of Metallic Zinc. 
 
 Heat a granule of zinc in an open crucible, noting and explain- 
 ing all changes which take place during the heating and subse- 
 quent cooling. Compare with the conduct of silver under sim- 
 ilar conditions. Tabulate the melting-points of the metals of 
 the first and second groups of the periodic system and formulate 
 a statement of the relation of melting-points to atomic weights. 
 
 96. The Equivalent Weight of Zinc. Quantitative. 
 
 Repeat the manipulation described in Exp. 93 using about 
 0.1 gram of zinc, and place in the test-tube a small piece of 
 platinum wire with which the zinc may come in contact, to 
 hasten its solution. Calculate the equivalent weight of zinc. 
 
100 MAGNESIUM, ZINC AND MERCURY 
 
 97. Preparation of Zinc Sulfate. 
 
 a. Make a suspension of 5 grams of finely powdered zinc 
 oxide in water, test this with litmus, and dissolve by adding 
 acetic acid, or dissolve in dilute hydrochloric acid and add 
 sodium acetate. Explain why these two procedures would 
 give the same hydrogen ion concentration. Place this solution 
 in an erlenmeyer flask fitted with a one hole stopper carrying a 
 delivery tube projecting below the surface of the liquid, and 
 pass in hydrogen sulfide with the stopper loose until the air in 
 the flask has been displaced, then close the flask, when hydrogen 
 sulfide will be supplied as fast as absorbed, and allow the flask 
 to stand until the solution is saturated. What is the precipi- 
 tate? Filter and wash, then dry the moist precipitate in an 
 evaporating dish, stirring constantly, and when entirely dry 
 heat in an open crucible to dull red for half an hour, or until 
 a small test portion shows the presence of appreciable amounts 
 of a soluble sulfate. Account for the formation of a sulfate. 
 Write the equation. Dissolve in water, warm and filter. Test 
 the filtrate with litmus and explain, then evaporate the solution 
 to crystallization. Devise and apply a test to determine 
 whether the insoluble residue is mostly zinc oxide or zinc sulfide. 
 
 6. Dissolve the zinc sulfate prepared in a, acidify with dilute 
 sulfuric acid and pass hydrogen sulfide into a portion of the 
 solution, then add an equal volume of sodium acetate solution 
 and again pass in hydrogen sulfide. Account for the difference. 
 
 c. To two other portions of the zinc sulfate solution add 
 sodium hydroxide till a permanent precipitate forms, then to one 
 add an excess of sodium hydroxide, and to the other hydro- 
 chloric or sulfuric acid. Is zinc hydroxide a base or an acid? 
 Indicate its products of ionization and the effect which the 
 presence of a base or an acid has upon its ionization. 
 
 98. The Preparation and Properties of Mercurous and Mercuric 
 Nitrate. 
 
 a. Clean 5 grams of mercury by pouring through a dry filter 
 paper through which a pin hole has been pricked at the apex. 
 
MERCUROUS AND MERCURIC NITRATES 101 
 
 Calculate the amount of concentrated nitric acid required to 
 convert this to mercurous nitrate (HgN0 3 ) and take a little less 
 (why?) than the calculated amount of the concentrated acid 
 (specific gravity 1.40,) and dilute with twice its volume of water. 
 Dissolve the mercury in this, warning gen.tly whe,n, necessary 
 to keep the reaction going, and after. -it 'has 'apparently ceased 
 (a good indication of complete reaction is:tJ6',fd7it^tioA t of a 
 yellow precipitate on diluting a small portion' of the solution' with 
 an equal volume of water explain this test after the experi- 
 ment has been completed) pour off the solution from the residue 
 of undissolved mercury, cool and allow to crystallize. To 
 separate portions of the mother liquor add hydrochloric and 
 sulfuric acids and ammonium hydroxide. Compare the solu- 
 bility of mercurous nitrate, sulfate and chloride. Add water 
 to some of the crystals, then warm and test the water solution 
 with litmus. Account for the test. What is the precipitate? 
 Is it soluble in water? In nitric acid? 
 
 b. Dissolve the rest of the mercurous nitrate in as little dilute 
 nitric acid as possible and warm, adding concentrated nitric acid 
 as necessary till brown fumes are no longer given off. What 
 reaction is taking place ? Calculate the amount of concentrated 
 nitric acid required to convert 5 grams of mercury to mercuric 
 nitrate (Hg(N0 3 ) 2 ) and compare with the quantity required for 
 mercurous nitrate. Crystallize this salt. How does its solu- 
 bility compare with that of mercurous nitrate? Dissolve some 
 of the crystals in water, adding dilute nitric acid if necessary, 
 and test the solution with hydrochloric and sulfuric acids and 
 ammonium hydroxide, and compare with the mercurous com- 
 pounds. Does the ammonium hydroxide form mercuric hydrox- 
 ide? Devise and apply a test to determine this. Add some 
 bright metallic copper to the mercuric nitrate solution. Where 
 is mercury in the electromotive series? What metals will be 
 precipitated from solution by mercury? Add a small amount 
 of zinc to mercuric nitrate solution and warm till the zinc is 
 nearly all dissolved, then test the solution by adding hydro- 
 chloric acid. What was the effect of the zinc? 
 
CHAPTER XIX 
 -'ALUMINIUM 
 
 99. Aluininates and Aluminium Hydroxide. 
 
 Dissolve about 2 grams of aluminium in dilute potassium 
 hydroxide, avoiding an excess of the alkali. What remains in 
 solution? What gas is evolved? Filter the solution and pass 
 hydrogen sulfide into a small portion. Is a sulfide precipitated? 
 Prove your answer experimentally. Pass carbon dioxide into 
 the main portion of the solution. What is the precipitate? 
 Filter and wash, and show experimentally that it is not a car- 
 bonate. Test small portions of this precipitate for solubility in 
 acids and alkalis, and explain in terms of the ionization theory. 
 Dissolve the main portion of the precipitate in hydrochloric 
 acid, evaporate until crystals appear on the surface of the 
 liquid, which should be kept strongly acid, and crystallize by 
 cooling and spontaneous evaporation. Dry the crystals, dis- 
 solve a small portion in water and pass in hydrogen sulfide. 
 Compare with the action of hydrogen sulfide on the alkaline 
 solution and explain. Heat the rest of the crystals in a dry 
 test-tube till decomposition is complete, noting the odor of the 
 gas given off. What is it? Is the residue a chloride? Test 
 this experimentally. 
 
 100. Water of Crystallization in Alum. Quantitative. 
 
 Calculate the quantities of aluminium sulfate (hydrated) and 
 potassium sulfate necessary to make 5 grams of the crystallized 
 aluminium alum, and dissolve these amounts in 20 c.c. warm 
 water, cool and allow to crystallize slowly. Dry the crystals on 
 filter paper, then by standing in the open air, and after they are 
 thoroughly dry determine the water of crystallization by 
 placing about 1 gram in a clean, dry crucible which has been 
 
 102 
 
ALUM 103 
 
 previously weighed, and weighing accurately. Cover this crucible 
 and place in another crucible of the same size and heat with a 
 flame not more than two inches high, and at least two inches from 
 the bottom of the lower crucible. After most of the water has 
 been driven off the lower crucible may be removed and the small 
 flame brought just below, but not quite touching the crucible 
 containing the alum, and the heating continued until constant 
 weight is attained. Too high a temperature will cause decom- 
 position of the aluminium sulfate. Calculate the per cent of 
 water in alum from this experiment, and from the formula, and 
 compare. 
 101. Aluminium in Kaolin. 
 
 Fuse some potassium pyrosulfate in a crucible and add pow- 
 dered kaolin in small portions until the mass will no longer 
 melt, then cool, dissolve in boiling water, filter and add am- 
 monium hydroxide to the filtrate. What is formed? Test 
 its solubility in potassium hydroxide and in hydrochloric 
 acid. 
 
CHAPTER XX 
 LEAD AND TIN 
 
 102. The Properties of Metallic Lead and of Lead Salts. 
 
 a. Test the solubility of metallic lead in hydrochloric, acetic, 
 sulfuric and nitric, acids. To solutions of lead nitrate add each 
 of the above acids, first cold, then warm. From the informa- 
 tion so gained explain the effects of the acids on metallic lead. 
 From the position of lead in the electromotive series would you 
 expect it to dissolve in these acids? 
 
 b. Add sodium hydroxide solution to lead nitrate solution, 
 first in small amount then in excess, and pass hydrogen sulfide 
 into the alkaline solution. Devise and apply a test to determine 
 whether the sulfide or hydroxide is precipitated here. Pass 
 hydrogen sulfide into a hot solution of lead acetate, acidified 
 with hydrochloric acid. Is the sulfide or hydroxide precipi- 
 tated? Compare these results with those obtained on similar 
 treatment of aluminium. Which element shows the more 
 basic properties? From the positions of lead and aluminium 
 in the Periodic System which would you expect to show most 
 basic properties? 
 
 103. The Oxides of Lead. 
 
 a. Heat a small sample of red lead in a dry test-tube, noting 
 the change in appearance, and at the same time test for oxygen. 
 After decomposition is complete divide the residue into two 
 portions; to one add dilute nitric acid and to the other sodium 
 hydroxide solution. Shake and filter if necessary, then pass 
 hydrogen sulfide into each solution. What has been the re- 
 action in each case? Is each color change on heating the red 
 lead accompanied by decomposition? Test this by heating 
 another small portion till only the first change has occurred, 
 then cooling. 
 
 6. Treat another small sample of red lead with dilute nitric 
 acid, stirring till the change is complete. Filter and add sodium 
 hydroxide solution to the filtrate, first in small amounts then 
 
 104 
 
TIN 105 
 
 in excess, and account for the changes. What is the residue 
 insoluble in nitric acid? Is it soluble in sodium hydroxide 
 solution? In hydrochloric acid (hood)? What are the prod- 
 ucts of the reaction with hydrochloric acid? Warm to dis- 
 solve the ] ead chloride and test the solution by adding sodium 
 Itydroxids solution and compare with the reaction of the nitrate 
 (the first filtrate, above). 
 104. Stannous and Stannic Tin. 
 
 a. Dissolve about 1 gram of tin in concentrated nitric acid 
 diluted with an equal volume of water, filter and wash. What 
 is the precipitate? Keep the filtrate and at the end of this 
 experiment devise and apply a test to determine the presence 
 or absence of tin. Dissolve a small portion of the precipitate 
 in potassium hydroxide solution. Dissolve the rest in as little 
 dilute hydrochloric acid as possible (A). What is formed in 
 each case? Test small portions of each of these solutions by 
 saturating with hydrogen sulfide, then acidify with hydrochloric 
 acid and account for every change that takes place in each solu- 
 tion. Dilute a small portion of the hydrochloric acid solution 
 (A) and add potassium permanganate solution (hydrochloric 
 acid may reduce potassium permanganate solution if sufficiently 
 concentrated). Is this stannous or stannic tin? 
 
 6. Dissolve 1 gram of tin by warming with concentrated 
 hydrochloric acid diluted with an equal volume of water, avoid- 
 ing an excess of the acid. Dilute and add potassium perman- 
 ganate solution to a small portion. Account for the effect. 
 Add potassium hydroxide solution to another portion till a 
 permanent precipitate forms, which on addition of more potas- 
 sium hydroxide and warming redissolves. What is in solution? 
 Pass hydrogen sulfide into this alkaline solution and also into 
 the acid solution. Compare with the same tests in a and ac- 
 count for the difference. 
 
 105. The Hydrogen Equivalent of Tin and Its Atomic Weight. 
 Quantitative. 
 
 a. Repeat the manipulation described in Exp. 93 using about 
 0.1 to 0.15 gram of tin, and place in the test-tube a small piece 
 
106 LEAD AND TIN 
 
 of platinum wire with which the tin may come in contact, to 
 hasten its solution. Dissolve the tin in concentrated hydro- 
 chloric acid instead of the dilute, warming if necessary. 
 Calculate the equivalent weight of tin, and by referring to 
 Exp. 52 calculate the equivalent weight of tin from that 
 experiment. Account for the difference. 
 
 b. Apply Dulong and Petit's Law to determine the atomic 
 weight from these equivalents by determining the specific heat 
 as follows. Take a piece of compact tin weighing 5 to 10 grams, 
 or if a compact piece is not available melt the required amount 
 of feathered or granulated tin in a crucible, and file down until 
 a bright smooth surface is obtained. Weigh this to within 
 0.1 gram, place in a dry test-tube, and immerse the test-tube in 
 a beaker of boiling water for ten to fifteen minutes, until the tin 
 shall have reached the temperature of the water. Measure out 
 in a small dry beaker 25 c.c. of water at the room temperature, 
 and provide it with a stirrer made by flattening the end of a 
 glass rod. Read the temperature of the water carefully, 
 remove the tin from the boiling water and drop it directly 
 from the test-tube into the beaker of water. Stir well, and after 
 a moment or two read the temperature of the water again. If 
 the temperature is read too soon error will result from incom- 
 plete distribution of the heat between the tin and the water, 
 if too late, there may be an appreciable error introduced by 
 loss of heat by radiation. Now read the temperature of the 
 boiling water and from this data calculate the specific heat of 
 the tin as follows: The change in temperature of the water 
 times its volume measured in c.c. gives roughly the number 
 of calories supplied by the tin (this is introducing a slight 
 error by ignoring the heat capacity of the beaker and stirrer, 
 and loss of heat by radiation), and that divided by the change 
 in temperature of the tin times its weight gives the amount of 
 heat supplied by 1 gram of the tin in changing its tempera- 
 ture by 1, which is its specific heat. Apply Dulong and 
 Petit's Law and give the true atomic weight of tin. 
 
CHAPTER XXI 
 CHROMIUM 
 
 106. The Preparation of Sodium Chromate from Chrome Iron 
 Ore. 
 
 Mix thoroughly 5 grams of sodium peroxide with 2 grams of 
 well ground chrome iron ore and heat gently, keeping the 
 mixture just melted for about twenty minutes. Cool and dis- 
 solve in a small amount of warm water and filter. What is the 
 precipitate? Add hydrochloric acid to the filtrate as long as the 
 orange color formed as the acid comes in contact with the 
 solution changes back to the original yellow color on stirring. 
 Explain these changes after the experiment is completed. 
 What two salts are now present in this solution? Look up 
 their solubilities and devise a method of taking advantage of 
 their differences in solubility in hot and cold water, to accom- 
 plish their separation, and carry this out. Dissolve some of the 
 crystals of sodium chromate and to one portion of the solution 
 add silver nitrate solution, then acidify with nitric acid. To 
 another portion add sulfuric acid and note the color change, 
 then warm, adding alcohol drop by drop, and again note the 
 color change. Compare these colors with those of chromates, 
 dichromates and any chromium salt, and account for each 
 change of color in the solution. 
 
 107. Chromyl Chloride. 
 
 Grind together 2 grams of potassium chromate and 1 gram of 
 sodium chloride, place a small portion in a dry test-tube and add 
 two or three drops of concentrated sulfuric acid slowly. Warm 
 and collect a drop or two of the distillate which condenses on 
 the walls of the test-tube by pouring it from the test-tube into 
 a beaker containing water. Test this solution with litmus. 
 
 107 
 
108 CHROMIUM 
 
 Add nitric acid and silver nitrate solution to one portion, to 
 another portion add silver nitrate solution without acidifying. 
 Explain the result in each case. 
 
 108. The Decomposition of Chromates by Heat. 
 
 a. Heat a small portion of well ground potassium dichromate 
 in a dry, hard glass test-tube. Do the crystals contain water of 
 crystallization? Is the substance volatile at this temperature? 
 Heat till reaction ceases (how may that point be determined?), 
 testing to determine what gas is evolved. Add water to the 
 residue in the test-tube after cooling and heat if necessary to 
 accomplish solution. Filter. What is the precipitate? Is 
 it soluble in sulfuric acid? What is the nitrate? 
 
 6. To 5 c.c. of saturated potassium dichromate solution add 
 concentrated sulfuric acid slowly till a permanent precipitate 
 forms. What is the precipitate? Wet a piece of filter paper 
 in this mixture and explain the color change. Dry the precipi- 
 tate by decanting as much of the liquid as possible and absorb- 
 ing the rest on a porous plate, and heat some of this dried 
 material in a dry test-tube. Is it volatile ? Does it decompose ? 
 Cool and treat the product with water and compare with the 
 products obtained by heating potassium dichromate. 
 
 109. Chrome Alum. 
 
 Calculate the amount of sulfuric acid necessary to make 
 chrome alum from 5 grams of potassium dichromate, reducing 
 with alcohol. Dissolve 5 grams of powdered potassium di- 
 chromate in a small amount of water, add the calculated amount 
 of sulfuric acid diluted with an equal volume of water, heat the 
 solution and add 50 per cent alcohol, drop by drop, till reduction 
 is complete. Set most of this solution aside to crystallize. To 
 a small portion add sodium hydroxide solution, first in small 
 quantities then in excess, and explain the changes. If chrome 
 iron ore were fused with sodium hydroxide instead of the 
 peroxide (Exp. 106) what would be the resulting compound? 
 
CHAPTER XXII 
 MANGANESE 
 
 110. Compounds Showing the Varying Valence of Manganese. 
 
 a. Melt 4 grams of potassium hydroxide in an iron crucible 
 and add gradually 2 grams of powdered manganese dioxide, 
 increasing the heat as necessary to keep the mixture just melted, 
 and stirring with an iron rod or file. In this procedure the 
 manganese undergoes auto-oxidation, part being oxidized to- 
 hexavalent manganese while the other part is reduced to- 
 trivalent. Write the equation, assuming that the oxides are 
 formed. Which of these oxides would be acidic, and what would 
 be the effect of the potassium hydroxide on the acid oxide? 
 Dissolve the fused mixture in distilled water and decant through 
 a filter. Keep this filtrate for d and the insoluble residue for b. 
 
 6. Wash the residue with distilled water until the wash water 
 is colorless. Some of the original manganese dioxide and some 
 manganic oxide will be present in this residue. Add dilute 
 sulfuric acid. Which oxide will dissolve? Allow to settle and 
 decant the clear sulfuric acid solution from the residue, (save 
 the residue for c), then make the solution alkaline with potas- 
 sium hydroxide. What is formed? Keep this test for com- 
 parison later. 
 
 c. Wash the residue insoluble in sulfuric acid (what is it?) 
 with distilled water till the wash water is colorless, then add 
 dilute sulfuric acid, warm and add alcohol or sodium sulfite 
 solution, drop by drop. What is the reaction? Make this solu- 
 tion alkaline with potassium hydroxide and compare with 
 manganous hydroxide, made by adding potassium hydroxide to 
 manganous sulfate, and with the manganic hydroxide prepared 
 above. Which is it? Pour off as much of the supernatant 
 liquid as possible and allow the precipitate to stand several 
 
 109 
 
110 MANGANESE 
 
 hours, and compare with the manganic hydroxide again. 
 Account for the change. 
 
 d. Pour a part of the green filtrate from the fusion mixture 
 into a large excess of water. What is formed? What is the 
 oxidizing agent? The reducing agent? Acidify another por- 
 tion with sulfuric acid and account for the change in color of the 
 solution and for the precipitate formed. Add sodium sulfite 
 solution to an alkaline and an acid solution of potassium per- 
 manganate, and to some of the green manganate solution. 
 Which oxide of manganese is formed when alkaline solutions 
 of manganates and permanganates are reduced? When an 
 acid solution of permanganate (can an acid solution of a man- 
 ganate be formed?) is reduced what is the valence of the man- 
 ganese in the resulting compound? Devise and apply a test to 
 prove your answer to these questions. 
 
 e. Add dilute nitric acid and some red lead to a 
 solution of mangarious sulfate, heat to boiling then allow the 
 precipitate to settle and note and account for the color of the 
 supernatant liquid. 
 
CHAPTER XXIII 
 IRON 
 
 111. The Preparation of Ammonium Ferrous Sulfate. 
 
 Dissolve 5 grams of iron nails or wire by heating with the 
 calculated amount of sulfuric acid diluted with four times its 
 volume of water (hood) adding water if necessary to replace that 
 lost by boiling, but keeping the final volume under 50 c.c. 
 When the reaction has ceased filter, add a few drops of dilute 
 sulfuric acid and the calculated amount of solid ammonium 
 sulfate necessary for ferrous ammonium sulfate, heat to boiling 
 (add water only if necessary to dissolve the ammonium sulfate 
 at the boiling temperature) and as soon as solution is complete 
 cool rapidly by holding under running water and stirring. 
 Drain the crystals from the supernatant liquid, wash once with 
 a small amount of cold distilled water, and dry the crystals on 
 filter paper and finally by standing in the air. Weigh and 
 calculate the per cent, yield (i.e. the percentage which the 
 amount formed is of the theoretical amount obtainable from 5 
 grams of iron). Dissolve a small portion and test the solution 
 with litmus. Save this preparation for Exps. 112 and 116. 
 
 112. Ammonium Iron Alum. 
 
 a. Dissolve 5 grams of the ferrous ammonium sulfate pre- 
 pared in Exp. Ill in 15 c.c. of distilled water, heat to boiling and 
 add concentrated nitric acid drop by drop with constant boiling 
 until the addition of more nitric acid does not cause the forma- 
 tion of a dark brown color (to what is this color due?). Crys- 
 tallize the alum by cooling and evaporation, and remove the 
 crystals from the liquid, dry on filter paper and place in a 
 stoppered bottle for use in Exp. 113. Dissolve a small portion 
 and test the solution with litmus. Compare with the ferrous 
 salt. 
 
 ill 
 
112 IRON 
 
 b. Test solutions of ferrous and ferric salts with ammonium 
 hydroxide, potassium permanganate solution and potassium 
 iodide solution. Contrast the effects of the two iron salts 
 in each case. Write equations for each raction. 
 
 113. Determination of the Water of Crystallization and of Ferric 
 Oxide in Ammonium Iron Alum. Quantitative. 
 
 Weigh a clean dry porcelain crucible, place in it 0.2 to 0.5 
 gram of the ammonium iron alum prepared in Exp. 112 (if 
 these crystals have become opaque, or formed a brown deposit 
 on the surface they will not give good results as decomposition 
 has already started) and determine the water of crystallization 
 according to the method used for aluminium alum, Exp. 100. 
 After constant weight has been attained at the low temperature 
 required for driving off the water, heat the crucible slowly to red 
 heat with the direct flame, cool and weigh, and repeat till it 
 comes to constant weight. The residue left in the crucible is 
 ferric oxide. From the weights of the water, the ferric oxide 
 and the original weight of the alum taken calculate the per cent, 
 of water and iron in the alum. 
 
 114. The Hydrolysis of Ferric Salts, and Colloidal Ferric 
 Hydroxide. 
 
 a. Heat a small crystal of hydrated ferric chloride in a dry, 
 soft-glass test-tube, testing the gas evolved with moist litmus 
 paper. What is evolved? Heat till the test-tube is entirely 
 dry, cool, add water, filter and test the filtrate for iron. What is 
 the insoluble residue? 
 
 b. Make 50 c.c. of a 5 per cent, solution of ferric chloride, add 
 sodium carbonate solution as long as the precipitate first formed 
 redissolves (explain the phenomena observed here) and if a 
 permanent precipitate forms add dilute hydrochloric acid till 
 it dissolves. Dialyse in a collodion sack (Exp. 71 6) changing 
 the water frequently at first and test the dialysate for chlorides. 
 When the dialysate shows only traces of chlorides after standing 
 in contact with the sack for about an hour remove the sack and 
 
FERROUS AMMONIUM SULFATE 113 
 
 test the solution of colloidal iron for chlorides. Add a concen- 
 trated solution of sodium chloride to this solution. Is this a 
 physical or a chemical change? Add a dilute acid to another 
 portion of the colloidal iron hydroxide solution and account for 
 the change in color. 
 
 c. Recall the effect of ferrous and ferric sulfate solutions on 
 litmus. In which valence does iron show most marked metal- 
 lic properties? 
 
 115. Cast Iron and Steel. 
 
 Dissolve small samples of cast iron and steel in dilute hydro- 
 chloric acid, testing the gases evolved for hydrogen sulfide. 
 Note the amount and appearance of the residue in each case. 
 Filter and mix the dried residues with copper oxide and heat 
 each separately in a hard glass test-tube, testing the evolved 
 gas for carbon dioxide. Dissolve about a gram of steel filings 
 by stirring with potassium cupric chloride solution (what is the 
 white precipitate that forms and redissolves?). Note the 
 appearance of the residue and test the gases evolved when it is 
 heated with copper oxide. What elements besides iron are 
 present in cast iron and steel? 
 
 116. The Oxidation of Ferrous Ammonium Sulfate by Potas- 
 sium Permanganate and by Potassium Dichromate. Quan- 
 titative. 
 
 Make 0.1 formular solutions of potassium permanganate and 
 potassium dichromate by weighing accurately the amount re- 
 quired for 100 c.c., placing this amount in the measuring 
 cylinder, dissolving in distilled water, then adding water to 
 make just 100 c.c. (be careful that none of the solution is lost). 
 The rest of this experiment must be performed in one laboratory 
 period. Make 100 c.c. of a 0.1 formular ammonium ferrous 
 sulfate solution, weighing out the salt prepared in Exp. Ill 
 accurately (remember that these crystals contain water) and 
 place exactly 50 c.c. of this solution in each of two beakers. 
 Add to each solution 5 c.c. of dilute sulfuric acid, heat to boiling 
 and to the first beaker add the prepared solution of potassium 
 
114 IRON 
 
 permanganate from a burette until the purple color of the per- 
 manganate just becomes permanent, and record the amount of 
 permanganate solution required. To the second beaker con- 
 taining ammonium ferrous sulfate heated to boiling add potas- 
 sium dichromate solution from a burette ; 1 c.c. at a time, 
 testing a drop of the solution after each addition by bringing 
 it in contact with a drop of potassium ferricyanide solution on a 
 white dish or test plate. As long as a blue precipitate or solu- 
 tion is formed on mixing these drops there is still some unchanged 
 ferrous sulfate in the solution, and potassium dichromate 
 solution should be added only until the test drop no longer gives 
 a blue color. Record the amount of potassium dichromate 
 solution required. How do these two oxidizing agents compare 
 in oxidizing power, or in the amount of oxidation accomplished 
 by equal molecular quantities of each of the salts? In order to 
 make a normal solution of each so that the oxygen available 
 for oxidations should be equivalent to 1 gram of hydrogen per 
 liter, how much of each salt would be required for a liter of 
 solution? 
 
INDEX 
 
 Acetylene, 79 
 
 Air, determination of oxygen in, 68; 
 dissolved in water, composition 
 of, 69 
 
 Alcohol, the oxidation products of, 
 79 
 
 Alum, 102; ammonium iron, 111 
 
 Aluminates, 102 
 
 Aluminium, hydroxide, 102; in Kao- 
 lin, 103 
 
 Ammonia, 63 ; weight of a liter of, 64 
 
 Ammonium, ferrous sulfate, 111; hy- 
 droxide, 64; iron alum, 111 
 
 Antimony, 75 
 
 Arsenic, 73 
 
 Balance, 14 
 
 Barium bicarbonate, solubility of, 
 
 96; hydroxide, solubility of, 95; 
 
 sulfate, solubility of, 96 
 Bismuth, 75 
 
 Boiling point as a test for purity, 16 
 Borates of the heavy metals, 84 
 Boric acid, preparation from borax, 
 
 83; anhydride, 83 
 Boron, preparation, 83 
 Bromine, preparation of, 49 
 Bunsen burner, 10 
 Burette, 11 
 
 Calcium, bicarbonate, solubility of, 
 96; carbonate, 95; chloride, 97; 
 hydroxide, solubility of, 95; sul- 
 fate, solubility of, 96 
 
 Carbon, 76; dioxide, 76; dioxide, 
 density of, 77; monoxide, 77 
 
 Catalysis, 21 
 
 Chlorine, density of, 38; preparation 
 of, 38; properties of, 41 
 
 Chromate, sodium, 107 
 
 Chromates, decomposition of, 108 
 
 Chrome, alum, 108; iron ore, 107 
 
 Chromyl chloride, 107 
 
 Colloidal, ferric hydroxide, 112; sil- 
 icic acid, 81 ; sulfur, 56 
 
 Combining weight, of copper, 29; 
 of tin, 67 
 
 Concentration, of a solution of acid 
 or base, 59; on the speed of a re- 
 action, effect of, 23 
 
 Copper, 92; combining weight of, 29 
 Cupric compounds, 92 
 Cuprous compounds, 92 
 
 Definite proportion, law of, 20 
 Diffusion of gases, rate of, 28 
 Dulong and Petit 's law, 106 
 
 Electromotive series, 93 
 Elements and compounds, 18 
 Equilibrium, 53 
 Ethylene, 79 
 
 Ferric, hydroxide, colloidal, 112; 
 
 salts, hydrolysis of, 112 
 Ferrous ammonium sulfate, 111; 
 
 oxidation of, 113 
 
 General directions, 1 
 
 Hydrochloric acid, 42 
 
 Hydrofluoric acid, 52 
 
 Hydrogen, preparation and proper- 
 ties, 26 
 
 Hydrogen peroxide, concentration 
 of, 35; preparation and properties 
 of, 35 
 
 Hydrogen sulfide, 56 
 
 Hard water, 77 
 
 Hydrolysis, 85 
 
 Iodine, 52 
 
 Indicators, 86; quantitative, 88 
 lonization, 44; 85 
 
 Iron alum, 111; water of crystalli- 
 zation and ferric oxide in, 112 
 Iron, cast, 113 
 
 Law of definite proportion, 20 
 Law of multiple proportion, 36 
 Lead, properties of, 104; oxides of, 
 
 104; salts, properties of, 104 
 Lime, 95 
 
 Magnesium, equivalent weight of, 
 
 98; carbonate, 99; oxide, 99 
 Manganese, varying valence of, 109 
 Measuring instruments, 11 
 Melting point a test for purity, 17 
 Mercuric nitrate, preparation of, 101 
 
 115 
 
116 
 
 INDEX 
 
 Mercurous nitrate, preparation of, 
 
 100 
 
 Methane, 78 
 Multiple proportion, law of, 36 
 
 Neutralization, 59 
 Nitrates, behavior on heating, 66 
 Nitric acid, 65; oxidation with, 67 
 Nitric oxide, 65 
 
 Nitrites, behavior on heating, 66 
 Nitrogen, preparation and proper- 
 ties, 63 
 Nitrous oxide, 66 
 
 Orthophosphates, the preparation 
 of primary, secondary and terti- 
 ary, 72 
 
 Oxygen, determination of, in the air, 
 68; preparation and properties of, 
 21; weight of a liter of, 23 
 
 Phosphates of sodium, 72 
 
 Phosphorus, oxides of, 71 
 
 Pipette, 11 
 
 Potassium, acid sulf ate, 62; bromate, 
 50; chlorate, 47; dichromate, oxi- 
 dation by, 113; dichromate, sol- 
 ubility of, 32; hypochlorite, 46; 
 nitrate, 90; per chlorate, prepara- 
 tion of, 47; perchlorate, purity 
 of, 48; permanganate, oxidation 
 by, 113 
 
 Primary orthophosphates, 72 
 
 Pure substances and mixtures, 16 
 
 Quantitative experiments, directions 
 for, 3 
 
 Rock salt, purification of, 45 
 
 Secondary orthophosphates, 72 
 
 Silicic acid, 81 
 
 Silicon tetrafluoride, 81 
 
 Silver, chloride emulsion, 94; oxide, 
 94 
 
 Soap, 78 
 
 Sodium, carbonate by the Solvay 
 process, 89; chloride, purification 
 of, 45; chromate from chrome iron 
 ore, 107; sulf ate, 89; thiosulfate, 
 preparation and properties of, 58 
 
 Solubility of salts in water, quant., 32 
 
 Solutions, supersaturated, 33 
 
 Standard conditions for gases, 5 
 
 Stannic tin, 105 
 
 Stannous tin, 105 
 
 Steel, 113 
 
 Strontium, bicarbonate, solubility 
 of, 96; hydroxide, solubility of, 
 95; sulf ate, solubility of, 96 
 
 Sulfur, chemical properties of, 56; 
 physical properties of, 55; dioxide, 
 57 
 
 Sulfuric acid, 59 
 
 Table, of barometric corrections, 6; of 
 densities and composition of water 
 solutions of acids and ammonia, 7; 
 of solubility and density of gases, 
 9; of solubility of salts, 8; of vapor 
 pressures, 6 
 
 Tertiary orthophosphates, 72 
 Tin, atomic weight of, 105; com- 
 bining weight of, 67, 105; stannous 
 and stannic, 105 
 
 Water, composition of, 31; hard, 77; 
 of crystallization in alum, 102; of 
 crystallization in iron alum, 112; 
 of hydration, 34 
 
 Waters, natural, 34 
 
 Zinc, equivalent weight of, 99; prop- 
 erties of, 99; sulf ate, preparation 
 of, 100 
 
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