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Companion Texts to 4:his Outline 
 
 INORGANIC CHEMISTRY 
 
 BY 
 
 ALEXANDER SMITH 
 
 780 + XVIIIpp. With 106 Figures 
 
 Octavo, $2.25 net 
 
 GENERAL CHEMISTRY 
 FOR COLLEGES 
 
 BY 
 ALEXANDER SMITH 
 
 531 +XIII pp. With 67 Figures 
 
 Octavo, $2.15 net 
 
 NEW YORK. THE CENTURY CO. 
 LONDON. GEO. BELL AND SONS 
 
A LABOEATORY OUTLINE OF 
 GENERAL CHEMISTRY 
 
 BY 
 
 ALEXANDER SMITH 
 , 
 
 PROFESSOR OF CHEMISTRY IN THE UNIVERSITY OF CHICAGO 
 
 jfourtb 3E>ttfon. 
 
 REVISED IN COLLABORATION WITH 
 
 WILLIAM J. HALE 
 
 ASSISTANT PROFESSOR OF CHEMISTRY IN THE 
 UNIVERSITY OF MICHIGAN 
 
 r> 
 
 TWENT Y-SIXTH 
 
 NEW YORK 
 THE CENTURY CO. 
 
 1910 
 
<r 
 
 COPYRIGHT, 1907, BY THE CENTURY Co. 
 
 COPYRIGHT, 1899, BY ALEXANDER SMITH 
 
 Set up and electrotyped June, 1899 
 
 Reprinted September, 1900 
 Second Edition, Revised, October, 1902 
 
 Reprinted August, 1905 
 
 Third Edition, Revised, June, 1907 
 
 Fourth Edition, Revised, January, 1908 
 
 Reprinted August, 1908 
 
 Reprinted April, 1909 
 
 Reprinted January, 1910 
 
 Stanbope press 
 
 H. OIL.60N COMPANY 
 BOSTON, U. 6. A. 
 
PREFACE TO THE THIRD EDITION 
 
 THE extensive use of the previous editions seems to have 
 shown the acceptability of the general plan of the book. This 
 conclusion has been confirmed by the fact that the second 
 edition has been translated into German, and that a transla- 
 tion of the present one into Russian is being prepared. In the 
 present, therefore, the fundamental features of the previous 
 edition have been preserved. In brief, the aim has been to 
 furnish the basis for a systematic, coherent, and instructive 
 study of the elements of chemistry from the modern stand- 
 point. 
 
 In the effort to make misapprehensions and mistakes as 
 nearly impossible as may be, the directions have been entirely 
 rewritten, and in many cases have been amplified, and a num- 
 ber of the experiments have been modified. An entirely new 
 set of figures has also been drawn. To render the exercises more 
 instructive, and still further to discourage mechanical work, a 
 larger number of questions has been inserted. With the same 
 end in view, data in regard to solubility have been introduced 
 (Appendix) as a new feature, and their use hi explaining 
 chemical phenomena has been illustrated in many experiments. 
 The rationalizing value of using the conceptions of chemical 
 dynamics, the electromotive series, and the degrees of ioniza- 
 tion has been emphasized by more frequent references. 
 
 Some of the formal quantitative experiments have been 
 modified, and the directions have been made clearer. The 
 value of these experiments has been found to lie chiefly in the 
 basis they give for clear understanding of the difficult subject 
 of combining, atomic, and molecular weights. 
 
 When quantitative experiments were first used in elementary 
 chemistry it was hoped that they would also assist in develop- 
 ing an abiding realization of the quantitativeness of all chem- 
 ical phenomena and, as a consequence, make all the thought 
 and work of the student more rigorous. In the experience of 
 the authors, however, quantitative experiments of the usual 
 kind fail to accomplish this important result. Students who 
 have performed such experiments still add a test-tube full of 
 sulphuric acid to a liquid known to contain only a trace of 
 a compound of lead, and still think less than a dozen bubbles 
 of hydrogen sulphide sufficient to precipitate the lead from 
 
 M88208 
 
vi PREFACE TO THE THIRD EDITION 
 
 10 cc. of approximately normal lead nitrate solution (see 74 g 
 and note 36, p. 67). They attempt to make potassium chlorate 
 without considering that a few bubbles of chlorine (perhaps 
 liberally mixed with air) will not saturate three grams of 
 potassium hydroxide (see 57 a), or they take too much water 
 and then, not having considered the solubilities and, therefore, 
 not knowing what is wrong, throw away the product and lose 
 valuable time by starting entirely ab initio. The failures 
 which result from this lack of a sense of quantity are innum- 
 erable, and the discouragement often a serious hindrance to 
 ultimate success. The fault, of course, is in the instruction, 
 and the remedy lies in exercises and questions devised to cul- 
 tivate this missing sense. It is to meet this situation that the 
 tables of solubilities have been introduced and have been 
 referred to frequently (see, e.g., 64, 65, 126 a, 127, 137 c-f, 139 d). 
 With the same object, the tables of degrees of ionization have 
 been utilized (see e.g., 64, 66), and the varying degrees of 
 activity of acids have been observed (e.g., 164) and measured 
 (120). Still further to cultivate rational experimentation, the 
 solutions on the side-shelf should be approximately normal 
 (or in simple multiples or submultiples of this concentration), 
 and the student may then be led to note the concentrations 
 and, in many experiments, to take suitable quantities accord- 
 ingly. The importance of the point of view indicated in the 
 foregoing can hardly be overestimated. Genuine success in 
 business or in the professions, and often even the mere making 
 of a livelihood, depend so largely on ability to reason quanti- 
 tatively that practice in this kind of reasoning is of inestimable 
 value in education. If, on the contrary, the work in chemistry 
 is purely haphazard in this respect, the study of the science 
 may easily be a positive detriment rather than a benefit, and 
 that whether the student ultimately makes direct use of his 
 knowledge of the science or not. 
 
 If it appears that these changes have made the work more 
 difficult, it must be remembered that valuable knowledge can 
 be obtained only by effort, and that the value of the knowledge 
 is in proportion to the effort, provided the latter is directed 
 rationally along instructive lines. Easy chemistry must be 
 superficial and empirical, in proportion to its simplicity. It 
 is easy to perform experiments mechanically; it is necessarily 
 more difficult to interpret the results and extract all that they 
 can teach. 
 
 The book is intended for beginners in colleges, universities, 
 and professional schools. It must be understood, however, 
 that no one class is expected to perform all the experiments. 
 
PREFACE TO THE THIRD EDITION vii 
 
 Only from one-half to three-quarters of the whole material can 
 be covered in thirty-three weeks, by a student working four to 
 six hours a week. The authors have found no difficulty in 
 arranging a course only twelve weeks long, and utilizing con- 
 siderably less than half the contents. The outline is subdivided 
 into numerous small paragraphs, so that each instructor may be 
 able to make such a selection as will suit the work he desires 
 to give. It is hoped that the whole body of material is suffi- 
 ciently great to permit the arrangement of a course of almost 
 any character. Thus, many or few quantitative experiments 
 may be given, and many or few theoretical matters illustrated. 
 Emphasis may be laid on work leading to analysis, or some 
 of that work may be sacrificed in order to include a larger 
 number of preparations. The student may even be directed 
 in certain paragraphs to ignore the questions. Finally, the 
 order of the experiments may be altered without serious dis- 
 turbance. 
 
 The recent great improvement in the work in chemistry hi 
 secondary schools makes it desirable to recognize and encour- 
 age this work by outlining a different selection of experiments 
 for those who have studied the science before, and, if possible, 
 to give such students separate class-room instruction. 
 
 Sample selections of experiments for beginners and for more 
 advanced students have been given in a separate pamphlet, 
 which will be sent to instructors, upon request, by -publishers. 
 This pamphlet includes also lists of apparatus and chemicals 
 and other data helpful in the organization of the laboratory 
 instruction. 
 
 In the preparation of this edition the authors have received 
 helpful suggestions from Prof. S. Lawrence Bigelow of the 
 University of Michigan, from Prof. Ralph H. McKee of Lake 
 Forest University, and from Prof. H. N. McCoy, Dr. C. M. 
 Carson, and Mr. T. B. Freas of the University of Chicago, 
 as well as from many others. We wish here gratefully to 
 acknowledge the improvements which the Outline owes to 
 these suggestions. 
 
 THE AUTHORS. 
 
 CHICAGO AND ANN ARBOB, 
 May, 1907 
 
PREFACE TO THE FOURTH EDITION 
 
 THE third revised edition having been exhausted within a few 
 months of its publication, the opportunity is taken to introduce 
 some needed alterations. Aside from a few corrections, only 
 one considerable change is made. This consists in the transfer 
 of Chapter V of the last edition bodily so that it follows Chapter 
 VII of the same edition. This places the quantitative experi- 
 ments on equivalent weights after the work on water, chlorine 
 and hydrogen chloride instead of before it. In the " Introduc- 
 tion to General Inorganic Chemistry," to which the Outline is a 
 companion volume, the discussions of molecular and atomic 
 weights and of the atomic hypothesis occur at this point and are 
 therefore now contemporaneous with the quantitative experi- 
 ments dealing in part with the same subjects. Thus there have 
 been adjusted the two difficulties in making a convenient scheme 
 of work, which were formerly occasioned by the lack of laboratory 
 work to accompany the molecular and atomic hypotheses and 
 their applications, and, at a different stage, the great excess of 
 laboratory work over class-room work at the time when the main 
 set of quantitative experiments was being performed. Many 
 instructors* will feel that, apart from this mere matter of peda- 
 gogical convenience, the postponement of the quantitative 
 work is in itself advantageous, because, after more laboratory 
 experience, greater ease of performance and greater accuracy in 
 results may now reasonably be expected. 
 
 This opportunity may be taken to call attention to the 
 change in the nomenclature of ionic substances (p. 55) in 
 Chapter X et seq. This considerable departure from the system 
 used in the text-book was made only after careful considera- 
 tion. The nomenclature now employed in the outline, although 
 it has been developed chiefly since the text-book was written, 
 has already come into fairly consistent use in American chemi- 
 cal literature. In Great Britain no one system has established 
 itself. In Germany a plan similar to that here adopted has been 
 used by Ostwald and others for several years. 
 
 THE AUTHORS. 
 
 November, 1907. 
 
 viii 
 
CONTENTS 
 
 CHAPXEBS PAGES 
 
 GENERAL INSTRUCTIONS 1 
 
 I. MANIPULATION , 3 
 
 II. CHARACTERISTICS OF CHEMICAL PHENOMENA. ... 9 
 
 III. OXYGEN 14 
 
 IV. HYDROGEN 19 
 
 V. WATER AND SOLUTION 23 
 
 VI. CHLORINE AND HYDROGEN CHLORIDE 29 
 
 VII. EQUIVALENT WEIGHTS, FORMULA, EQUATIONS ... 33 
 VIII. BROMINE, IODINE, FLUORINE, AND THEIR COMPOUNDS 
 
 WITH HYDROGEN 41 
 
 IX. DOUBLE DECOMPOSITION. OXYGEN COMPOUNDS OF 
 
 THE HALOGENS. HYDROGEN PEROXIDE .... 48 
 X. loNlZATION AND INTERACTIONS OF ACIDS, BASES, 
 
 AND SALTS 55 
 
 XI. SULPHUR 65 
 
 XII. THE ATMOSPHERE, NITROGEN, AMMONIA ..... 76 
 
 XIII. OXIDES AND OXYGEN ACIDS OF NITROGEN .... 81 
 
 XIV. PHOSPHORUS ' . 87 
 
 XV. CARBON . 90 
 
 XVI. THE ACTIVITY OF ACIDS MEASURED CHEMICALLY . 96 
 
 XVII. SILICON AND BORON 98 
 
 XVIII. METALLIC ELEMENTS OF THE ALKALIES 100 
 
 XIX. METALLIC ELEMENTS OF THE ALKALINE EARTHS. . 107 
 
 XX. COPPER AND SILVER Ill 
 
 XXI. MAGNESIUM, ZINC, CADMIUM, MERCURY 116 
 
 XXII. ALUMINIUM, TIN, LEAD 120 
 
 XXIII. ARSENIC, ANTIMONY, BISMUTH 123 
 
 XXIV. CHROMIUM, MANGANESE 126 
 
 XXV, IRON, COBALT, NICKEL 129 
 
 APPENDIX . 132 
 
GENEKAL INSTRUCTIONS. 
 
 NOTES 1-17. 
 
 Read the "Regulations" posted in the laboratory. Read also, 
 attentively, the following notes: 
 
 Note 1. Provide yourself with a note-book and make a careful 
 permanent record immediately after each experiment. Enter the 
 numbers and titles of the paragraphs of the outline systematically. 
 State (1) what you did, if anything beyond the directions, but do 
 not copy the printed directions themselves, (2) what you observed, 
 (3) what conclusions you drew. A sketch of the apparatus will 
 enable you to recall the circumstances of the experiment, if later 
 reference to it is necessary. This note-book, when called for, is 
 to be handed in for inspection. 
 
 The blank pages in this Outline are not intended for the final 
 notes. They may be used for individual suggestions given by 
 the instructor, preliminary notes, record of weighings, etc. 
 
 The directions have been expressed with the utmost care and 
 brevity. Every word is significant. Italics are therefore nowhere 
 employed for the purpose of emphasis. 
 
 Note 2. Whenever an interrogation point or a direct ques- 
 tion appears a corresponding note should appear in the note-book. 
 The " (?)" indicates something to be observed and recorded. 
 
 Note 3. The very numerous questions asked in the course of 
 this outline are intended to be answered, not by speculation, but 
 by careful observation and reasoning based on the results of this. 
 Very often the student will find it necessary to devise and carry 
 out further experiments of his own before a satisfactory answer is 
 obtained. When a question occurs to you, endeavor by reflection 
 and study to answer it yourself before consulting an instructor. 
 
 Note 4. In many cases the work outlined could not in itself 
 furnish the basis for an answer, and fuller investigation of the 
 point would require work beyond the time or ability at the dis- 
 posal of the beginner. Such questions are distinguished by an [R], 
 indicating that Reference to some authority (lecture, book, or 
 assistant) must be made. The number following the R is that of 
 the page in Alexander Smith's Introduction to General Inorganic 
 Chemistry, where the necessary information may be obtained. 
 The authority should be consulted, however, only after the experi- 
 ments have been made and the notes written up as far as possible. 
 Note 5. When a chemical change has been observed the 
 equation should always be given in the notes, but an equation 
 alone is never a sufficient record. 
 
 Note 6. Where the word [Instructions] appears, consult the 
 instructor before going further. 
 
 Note 7. In quantitative experiments, marked [Quant.], use the 
 finer balance, in all other cases the rough scales in the laboratory, 
 
 Note 8. The expression [Storeroom] indicates that the 
 sary apparatus is not included in the individual outfits. 
 
 1 
 
2 NOTES 1-17 
 
 Note 9. When the word [Hood] appears, the operation is not 
 to be conducted at the desk in the open laboratory. The appara- 
 tus is to be at once transferred to the hood provided for operations 
 involving ill-smelling gases or vapors. 
 
 Note 10. Where exact quantities are not indicated, very small 
 amounts of solutions (1 c.c. or less) should be taken. This advice 
 is given, partly to secure saving of material, but chiefly to avoid the 
 waste of time which working with large quantities always entails. 
 
 Note 11. To obtain the necessary chemical substances, do 
 not carry the bottles from the side-shelf to the desk. Bring a clean 
 test-tube for liquids and a watch-glass for solids. For the latter, 
 a piece of the paper, provided near the side-shelf, may also be 
 used. When too much of any reagent has been taken, do not 
 return it to the bottle. 
 
 Note 12. The chemicals are divided into two sets, each 
 arranged alphabetically according to the scientific names.* The 
 first set consists of solids in small bottles, the second of liquids. 
 The bottles and their places are numbered consecutively to facili- 
 tate accurate replacement, and scrupulous care must be takeft not 
 to disarrange them. Read the labels attentively, as there are 
 frequently several kinds of the same substance (e.g., pure, and 
 commercial, dilute, concentrated, and normal). 
 
 All materials are supplied through the storeroom service. Do 
 not therefore take side-shelf bottles, when found empty, to the 
 instructor, but to the storekeeper for refilling. 
 
 Note 13. The expression [From Instructor], however, indi- 
 cates one of a few special -substances for which the student must 
 apply to an instructor. 
 
 Note 14. The nine bottles on the desk contain aqueous 
 solutions of sodium hydroxide, sodium carbonate, and ammonium 
 hydroxide, which are not to be found on the side-shelves, and 
 sulphuric, hydrochloric, and nitric acids in concentrated and dilute 
 form. These acids are all commercial. The corresponding pure 
 concentrated and dilute acids will be found on the side-shelf, but 
 are to be used only when the outline so directs. 
 
 Note 15. When any acid gets upon the clothing, apply 
 ammonium hydroxide solution [On desk] at once. 
 
 Note 16. Burns, whether caused by contact with hot objects, 
 by acids, or by corrosive liquids like bromine, are rubbed gently 
 with a paste of sodium-hydrogen carbonate and water. All burns, 
 save the slightest, must afterwards be dressed with an aqueous 
 solution of boric acid (half -saturated) [Side-shelf] to prevent infec- 
 tion. Obtain the assistance of an instructor. 
 
 Cuts must be washed in running water and dressed with boric 
 acid as above, or with lanolin containing 2 per cent of boric acid. 
 
 Note 17. All students work independently, except where 
 cooperation of two students is expressly directed. 
 
 * A pamphlet containing lists of the chemicals and apparatus re- 
 quired will be furnished, on application, by the publishers. 
 
LABORATORY OUTLINE 
 
 CHAPTER^. 
 
 MANIPULATION. 
 
 1. The Outfit of Apparatus. As articles missing or found 
 imperfect when the course is completed will be charged for, 
 check the outfit of apparatus carefully by comparison with the 
 list. To do this, put all the articles on the top of the desk and 
 make a mark on the margin of the list opposite to the names 
 of such as you are able to identify, at the same time returning 
 the article to the cupboard or drawer. With the assistance of 
 an instructor, the remaining, unfamiliar articles can then be 
 checked also. 
 
 2. Instructions. 
 
 a. Read the general instructions and notes preceding this 
 chapter very carefully, and do not fail to observe them. 
 
 b. The number of blast-lamps and balances being limited, 
 the whole class cannot perform the experiments in this chapter 
 simultaneously in the order given. The order is, in any case, 
 a matter of indifference. Two students from the group under 
 each assistant will begin with glass-working (4) or weighing 
 (6 and 7 a and 6). The other students will meanwhile per- 
 form 3, 6, and 7 d at the desks. 
 
 c. Record in the note-book the results of 3, 6, and 7 only. 
 
 3. Bunsen Burner. 
 
 a. Attach the Bunsen burner by means of rubber tubing 
 to a gas connection, close the air-holes at the base, and light. 
 Now open the air-holes gradually and note the effect upon the 
 flame (?). What is the proximate cause of the difference in the 
 two flames ? When using the burner for heating purposes, 
 always regulate the air supply so as to get a noiseless, non- 
 luminous flame. 
 
 b. Determine the structure of each kind of flame. Which 
 parts are relatively hotter, and which cooler? This may be 
 ascertained by observing the quality of the glow produced in a 
 platinum wire held across the flame in several positions. A 
 match-stick quickly inserted may also indicate the cooler portions. 
 Make sketches showing the real form of the flame (?). Where 
 should you hold an object in the non-luminous flame, in order 
 
 8 
 
4 MANIPULATION [ 4 
 
 to get the greatest heating effect? Prove the presence of un- 
 burnt gas in the inner cone by inserting therein one end of a 
 narrow glass tube, and igniting the gas that issues from its 
 uppers end. . Which region is deficient in oxygen and which has 
 an excess? . cThe <f6rmer is called the reducing, the latter the 
 oxidizing region. 
 
 5. 'l^ringfroin.the- side-shelf, in a watch-glass [Note 12, p. 2], 
 & small J quantity of borax 3* (sodium tetraborate). Heat your 
 platinum wire, which has been inserted into the fused end of a 
 glass rod, and immerse the glowing end of the wire in the borax. 
 Use the wire in straight condition, without any loop. Now, 
 hold the wire in the flame, observe the behavior of the borax, 
 and explain [R 528. See Note 4, p. 1]. The bead must be small 
 to avoid its dropping off. 
 
 d. Bring the hot borax bead in contact with a minute particle 
 of manganese dioxide [Notes 11 and 12], heat in the flame near 
 the outer edge until the particle has dissolved, and observe the 
 color of the bead when cold. If the bead is opaque, too much 
 of the dioxide has been taken: throw the molten bead off and 
 start again. 
 
 e. Cut off the gas supply until the flame is about 6 cm. in 
 height. Close the air-holes until a luminous point appears 
 at the apex of the inner cone, and hold the bead containing the 
 manganese dioxide steadily in this luminous (reducing) portion. 
 Before withdrawing the bead, lower it into the inner cone of 
 unburnt gases to cool. Observe the color of the bead (?). 
 
 /. Reheat the bead in the oxidizing part of the flame (?). 
 4, Glass -Working [Blast-Lamp Table. Instructions, and 
 Note 18, below]. 
 
 a. Cut a small piece off the wide soft glass tubing and make 
 a test-tube out of it. 
 
 b. Make a test-tube of hard (? [R 607]) glass. 
 
 c. Connect two pieces of narrow glass tubing to make a 
 longer piece. 
 
 d. Make a T-tube by connecting two pieces of narrow glass 
 tubing at right angles to each other. 
 
 Note 18. To cut narrow glass tubing make a slight, trans- 
 verse scratch with the triangular file. Hold the tubing so that 
 the points of the thumbs are together opposite to the scratch, and 
 press forward with the thumbs so as to bend the tube away from 
 this mark. Wide tubing may be cut by making a deep scratch 
 completely round the tube and starting the crack by touching with 
 
 * In this Outline the common, or popular names are often used 
 intentionally. The systematic names must, in such cases, be found 
 by the student [B]. 
 
5] MANIPULATION 5 
 
 the red-hot end of a glass rod. Other methods may be shown by 
 the instructor. 
 
 Always round off (fire-polish) the sharp edges of freshly cut 
 tubes by softening in the Bunsen flame (why do the edges become 
 rounded?). In the case of test-tubes, and other tubes of wide bore 
 in which corks are to be inserted, the mouth must be strengthened 
 and rendered fit for receiving the cork. To accomplish this, heat 
 the edge in the flame and spread it slightly, but uniformly, by 
 rotating in it a pointed piece of charcoal, or by turning outward 
 the softened edge with the reverse end of a file. 
 
 In making test-oubes and in connecting pieces of tubing, distend 
 the softened parts by blowing immediately before allowing to cool, 
 otherwise cracks are likely to appear. 
 
 To bend glass tubing, never use the Bunsen flame (why ? 
 Note 3). Always employ an ordinary, flat, luminous flame of 
 fish-tail form. Hold the tubing lengthwise in the flame, not across 
 it (why?), turning the tube slowly round its axis to receive uniform 
 heating. Keep the tube straight until it is soft enough to bend by 
 its own weight. Finally, do not actually make the bend while the 
 tube is in the flame, but after removal from it (why?). 
 
 Hold a piece of tubing in the Bunsen flame, without rotating the 
 tube, and bend it while it is in the flame (?). Compare the bend 
 with one made in the proper way and account for the difference. 
 
 6. Construction of a Wash-Bottle. Select a good cork which 
 will fit the mouth of the largest flask and soften it by rolling 
 under the foot upon the floor while pressure is cautiously ap- 
 plied. Bore two parallel holes with 
 the cork-borer, and smooth them 
 by means of a rat-tail file [Note 19, 
 below]. Bend two pieces of glass 
 tubing as indicated in Fig. 1 [Note 
 18], fire-polish their edges and insert 
 them in the openings in the cork. 
 Make the nozzle by softening a 
 piece of glass tubing in the Bunsen 
 flame, drawing it to capillary dimen- 
 sions after removal from the flame, 
 cutting (6-8 cm. long), and fire- 
 polishing. Connect the nozzle by 
 means of a short piece of rubber Fig. x 
 
 tubing. Test the apparatus to see 
 
 that it is air-tight [Note 20]. Finally, fill the flask with distilled 
 water [Note 21]. 
 
 Note 19. The cork borer is usually made of brass, and the 
 edge is easily turned. Form the habit of examining the edge and 
 freshening it by cautious application of a triangular file [Instruc- 
 tions. Note 6] every time it is to be used. Do not hold the cork 
 
6 MANIPULATION [ 6 
 
 against the table while boring, as the edge of the tool may be 
 ruined. Hold the cork in the hand and bore from the narrow end 
 with care, exactly parallel to the axis. If the cork and borer are 
 rotated round their axes and the edge is fresh, very little force will 
 be required. The borer is purposely chosen so as to be smaller 
 than the tubing. Its use thus 1 permits the enlarging and smoothing 
 of the hole with the rat-tail file until a perfectly fitting bore has 
 been made. 
 
 Note 20. Before use, every piece of closed apparatus employed 
 in this and all succeeding experiments must be tested for air- 
 tightness, and rendered perfectly air-tight. In this instance place 
 in the flask enough water to cover the lower end of the longer 
 tube and transfer the rubber connection to the shorter glass tube 
 and close it with a clamp. Now, blow through the longer tube so 
 that a few bubbles of air pass into the flask. If the apparatus is 
 air-tight the water will rise in this tube when the mouth is with- 
 drawn and will remain in an elevated position. If the water grad- 
 ually sinks to its former level, the apparatus is not air-tight. 
 Examination of the holes in the cork may show defects, which can 
 be remedied only by boring a fresh cork more carefully. If the 
 cork itself contains pores, these may often be closed by thorough 
 wetting. Do not on any account employ paraffin or sealing-wax 
 to patch defective places in a cork ; use a fresh one. 
 
 Note 21. Distilled water is to be used for making solutions 
 and for rinsing glassware (why?). Common water is to be used 
 for all other purposes. 
 
 6. Use of the Simple Balance [Instructions. Quant.]. By 
 turning the screw attachment in front of balance case, release 
 the beam and pans, allowing the beam to swing. Observe the 
 excursions made by the pointer. Divide by two the total divi- 
 sions covered by the pointer in one full swing, and count off 
 from either end of the swing the divisions which this number 
 designates, thus finding the position of the true zero point. The 
 beam must swing freely during the observation : the zero is never 
 to be read with the beam at rest. 
 
 This observed zero point may lie a little to the right or to the 
 left of the marked zero. Note down its distance, in scale divi- 
 sions, from the marked zero. If it lies to the right, prefix to 
 the number of divisions the plus ( + ) sign; if to the left, the 
 minus ( ) sign. The zero of any one balance changes, and 
 must be redetermined every time a weighing is made. 
 
 Place a 10 g. weight in each pan [Note 22, below], and deter- 
 mine the zero as before. Add the .01 g. weight to the right- 
 hand pan, and find the reading about which the pointer now 
 oscillates. The difference in reading between this point and the 
 last determined zero point gives the deflection due to the .01 g. 
 weight. It may be used for estimating weights less than .01 g. 
 
7] 
 
 MANIPULATION 
 
 Note 22. Great care must be taken in the use of the balance 
 and weights. The pans of the former must be let down upon 
 their supports when not in use and every time weights or other 
 objects are to be added or removed. All objects to be placed upon 
 the pans must previously be carefully cleaned and dried. Solids 
 are placed upon a watch-glass or upon a piece of glazed paper, and 
 never directly upon the pans. The weights must be lifted from 
 their case by means of forceps, never by the fingers. They are 
 usually placed on the right-hand pan, the objects to be weighed 
 on the left. 
 
 Note 23. In reckoning results, count first by the places vacant 
 in the box, and check by count- 
 ing the weights themselves. This 
 will enable you to avoid the com- 
 monest error in weighing. Finally 
 record the weights in the note- 
 book, or laboratory outline, and 
 never upon loose sheets of paper, 
 as loss of the latter will necessi- 
 tate a repetition of the entire 
 experiment. 
 
 7. Measuring Vessels 
 [Quant.]. 
 
 a. Fit a burette (Fig. 2) 
 with a short piece of rubber 
 tubing and glass nozzle (see 6). 
 The withdrawal of liquid from 
 the burette is regulated by a 
 pinch clamp upon this rubber 
 connection, or by placing a 
 small piece of glass rod (with 
 fire-polished edges) in the 
 middle of the rubber tubing 
 to choke the bore. In this 
 latter case, pinching the tub- 
 ing surrounding the glass rod 
 will permit the liquid to flow 
 at any desired rate from the 
 nozzle. 
 
 b. Support the burette upon 
 a ring-stand by means of a 
 
 clamp. Now fill the burette Fig. a 
 
 with distilled water, drawing 
 
 off a portion to insure the complete removal of air from the 
 rubber tubing and nozzle. The last bubble of air may be re- 
 moved by turning the nozzle upward while the water is allowed 
 
8 MANIPULATION [ 7 
 
 to flow. Read the height of water by observing the lower 
 side of the meniscus and estimate to tenths of a division. 
 
 Clean and dry a small beaker carefully and weigh it [Quant.]. 
 Allow 10-20 c.c. of the distilled water to run from the burette 
 into the beaker, read the new level of the water, and ascertain 
 the volume taken by subtracting the readings. Weigh the 
 beaker again and ascertain the weight of the water by subtrac- 
 tion. Calculate from your figures the weight of 1 c.c. of water. 
 Keep the beaker and contents for use in c. 
 
 c. Place some distilled water in the graduated cylinder 
 and read its volume. Pour about 10 c.c. of this water into the 
 beaker (already partly filled with water) , read the volume of "the 
 remainder, and find by subtraction the volume poured out. 
 Now weigh the beaker once more and find the weight of this 
 water by subtracting the previous weight. Calculate the weight 
 of 1 c.c. of water. 
 
 Is measuring volume by burette or by cylinder more accurate 
 (1 c.c. of water at 4 C. weighs 1 g.), and why? 
 
 d. Measure by means of the cylinder, roughly, the volumes of 
 water your flasks and beakers hold, and record the figures. Fill 
 the vessels to a convenient height for use, and not to the brim. 
 
CHAPTER II. 
 
 CHARACTERISTICS OF CHEMICAL PHENOMENA. 
 
 8. Qualitative Study of Chemical Phenomena. 
 
 a. Rub a pinch (about 0.5 g.) of powdered roll sulphur with 
 a very small globule of mercury (use the dropper) in a mor- 
 tar (?). When the mass has been intimately ground together 
 for about five or ten minutes and no particles of mercury are 
 longer observable, place it in a dry test-tube and add about 
 3 c.c. of carbon disulphide [CARE. Keep away from flames], a 
 solvent for the free sulphur. Shake well (do not apply heat) 
 and then pour off the clear solution into the sink [Hood. Note 
 25, below]. Repeat, shaking well and pouring away the clear 
 liquid. Continue this operation until a sample of the decanted 
 liquid fails to give any sulphur by spontaneous evaporation 
 upon a watch-glass. What is the product remaining in the tube 
 [R 657]? By means of a few c.c. of carbon disulphide wash this 
 product out upon a filter paper [Note 24, below] and allow it to 
 dry. To show the composi- 
 tion of this substance place 
 it in a small dry test-tube 
 and heat strongly over a 
 flame. What odor is noticed 
 at the mouth of the tube? 
 What product is found to 
 sublime upon the walls of 
 the test-tube? Rub the sub- 
 limate with a glass rod (?). 
 What characteristics of chem- 
 ical change have you noticed 
 in the above experiment? 
 
 Note 24. Cut from a sheet 
 of filter paper a piece 7 cm. 
 (3 inches) square. Fold the 
 square twice in directions at 
 right angles to one another, so 
 as to obtain a square of one- 
 fourth the previous area. Cut 
 
 Fig. 3 
 
 off the loose corners with scissors so that a quadrant is formed. 
 Open the folded paper so that a cone is produced (Fig. 3), and 
 place the cone in a dry glass funnel. Smaller or larger pieces of 
 paper are taken according to the amount of the precipitate or 
 
10 
 
 CHEMICAL PHENOMENA 
 
 [8 
 
 of the liquid to be filtered, the smallest size that will serve the 
 purpose being preferred. The funnel must be chosen so that the 
 filter paper, after being trimmed and placed in position, does not 
 quite reach the edge, much less project above it. While in use, 
 Ene funnel is placed in a ring on the stand, and is not to be held 
 in the hand. 
 
 Note 25. Pour away all ill-smelling substances, like carbon 
 disulphide, in the sink in the hood and not in the ordinary sinks 
 or jars. 
 
 b. Gunpowder is made from saltpeter (potassium nitrate), 
 roll sulphur, and charcoal. Bring specimens of these three sub- 
 stances on watch-glasses from the side-shelf, and examine them 
 with respect to properties which can be used for recognition and 
 separation [R 36, 37, 39], such as appearance and solubility in 
 various solvents. Try the solubility of each in distilled water 
 and in carbon disulphide, using in the latter case thoroughly 
 dried test-tubes [Notes 25 and 26]. Do not judge of solubility 
 by the eye, but filter the mixture [Note 24], catch a few drops 
 of the liquid on a watch-glass, evaporate, and see whether there 
 is any greater stain on the glass 
 than the pure solvent would itself 
 have left. The taste (?) of the salt- 
 peter is a characteristic property. 
 
 Now place about 1 g. of gunpowder 
 in a large test-tube and add 5-10 
 c.c. of water. Shake well (after clos- 
 ing the test-tube with the thumb), 
 'warm gently, and filter [Note 24]. 
 Evaporate the filtrate upon a water 
 bath or over a beaker of boiling 
 water (Fig. 4). Describe and name 
 the residue. Dry (why ?) the filter 
 paper and its black contents over a 
 radiator or in a drying oven. Shake 
 the dried product with cold carbon 
 disulphide in a dry [Note 26] test- 
 tube, filter, and allow the filtrate to 
 evaporate [Hood] spontaneously (?) . 
 What remains upon the paper? Did 
 any chemical change occur during the manufacture of gun- 
 powder? 
 
 Why are the ingredients of gunpowder pulverized so finely and 
 mixed so intimately? 
 
 Note 26. To dry test-tubes or flasks quickly after recent 
 washing, warm them slightly in the Bunsen flame, connect with 
 
 Fig. 4 
 
9. 
 
 Dish No. 1. Dish No. 2. 
 
 Wt. of dish with lead (silver) 
 
 Wt. of dish, empty 
 
 Wt. of lead (silver) 
 
 Wt. of dish with lead (silver) chloride 
 
 Wt. of dish with lead (silver) 
 Wt. of chlorine 
 
 Wt. of lead (silver) 
 
 Wt. of chlorine 
 
 Wt. of lead (silver) chloride 
 
 No. 1. No. 2. 
 Per cent lead (silver) (z) 
 Per cent chlorine x' 
 
9] CHEMICAL PHENOMENA 11 
 
 the air supply of a blast-lamp a piece of glass tubing long enough 
 to reach to the bottom of the vessel, and allow a gentle stream of 
 air to flow through the tube and vessel. 
 
 9. Law of Definite Proportions * [Quant.]. 
 
 a. Obtain two pieces of pure lead of different sizes, about 
 0.5 g. and 0.6 g. respectively. Clean and dry two evaporating 
 dishes, taking care to remove any paper labels which may be 
 attached to them, and weigh each with care [Quant., balance]. 
 Place one piece of the lead in each and weigh again. The dif- 
 ference will give the exact weight of the lead. Be careful in 
 your notes, and in handling, to distinguish the dishes from one 
 another. 
 
 Dilute 7 c.c. of concentrated nitric acid [Desk] with 14 c.c. of 
 water in the graduated cylinder and add to the lead in each 
 dish 10 c.c. of the diluted (1 : 2) nitric acid. Cover each dish 
 with a watch-glass, convex side downward, and set upon the 
 water bath. When the action is over, and the lead has entirely 
 disappeared, rinse the lower surface of each watch-glass into the 
 dish by means of a little water from the wash-bottle. Then add 
 to the contents of each 5 c.c. of pure [Side-shelf] dilute hydro- 
 chloric acid. This causes the precipitation of a compound of 
 lead and chlorine whose weight is next to be determined. To 
 accomplish this, the water and other substances present (ex- 
 cept the lead chloride) being all volatile, the mixture is dried 
 by evaporation on a water bath. An extra Bunsen burner 
 [Storeroom] and two beakers containing water may be used as 
 baths, as in Fig. 4 [Hood]. A match inserted between the dish 
 and the beaker permits escape of the steam and prevents the 
 possible upsetting of the dish by the vapor. When the con- 
 tents of the dishes are dry, moisten the contents of each with 
 2 c.c. of pure [Side-shelf], concentrated hydrochloric acid and 
 dry once more. Now transfer the dishes, one at a time, to the 
 clay triangle supported on the ring-stand, and heat the whole 
 residue cautiously with a small Bunsen flame. Watch it nar- 
 rowly and do not allow any of it to melt. When the dishes are 
 cold, weigh each again with care. The increase in weight over 
 the previous value in each case gives the weight of chlorine 
 which has combined with the known weight of lead. 
 
 * Before beginning, always endeavor to ascertain the object of each 
 experiment. By this means confused work and much waste of time 
 will often be avoided, and significant facts to be observed and recorded 
 will not escape notice. Consider, first, carefully the title, as it will 
 usually indicate the object of the exercise. If the title is not at once 
 fully understood, look up the topic in a reference book before going 
 further. 
 
12 CHEMICAL PHENOMENA [ 9 
 
 Calculate from each of the two sets of data the percentage of 
 lead (x) and of chlorine (x'} in lead chloride : 
 
 Wt. of lead used: Wt. of lead chloride : : x : 100. 
 Wt. of chlorine found : Wt. of lead chloride : : x' : 100. 
 Compare the results of the two measurements and interpret. 
 
 b. To secure almost ideal results, two pieces of pure silver 
 foil of about 0.5 g. and 0.6 g. should be substituted for the lead 
 in a. 
 
 Dilute 5 c.c. of concentrated nitric acid [Desk] with 5 c.c. of 
 distilled water in the graduated cylinder and add to the silver 
 in each dish 5 c.c. of the diluted (1 : 1) acid. When the action 
 is over, and the cover-glasses have been rinsed, as in a, add 2 c.c. 
 of pure [Side-shelf] dilute hydrochloric acid to the contents of 
 each dish. This causes precipitation of a compound of silver 
 and chlorine. Evaporate the liquids to dryness [Hood] and 
 heat the residues at once until signs of melting are seen. 
 Weigh each dish. The increase in weight over the previous 
 value in each case gives the amount of chlorine which has com- 
 bined with the known weight of silver. 
 
 Calculate from each of the two sets of data the percentages 
 of silver (x) and of chlorine (x'} in silver chloride : 
 
 WU of silver used : Wt. of silver chloride : : x : 100. 
 Wt. of chlorine found : Wt. of silver chloride : : x' : 100. 
 
 Compare the results of the two measurements and interpret. 
 
 Before cleaning the dishes, transfer the silver chloride to the 
 bottle for silver waste. 
 
 c. Clean and dry two evaporating dishes, taking care to 
 remove paper labels which may be pasted upon them. Weigh 
 each dish [Quant., balance]. Place in each some sodium- 
 hydrogen carbonate, in the one about 1 g. and in the other about 
 2 g., and weigh again. Treat both alike, as follows: Dissolve 
 the solid in pure dilute hydrochloric acid [R 93], adding little 
 by little and covering with a watch-glass between successive 
 additions to avoid loss by spirting. When the solid has wholly 
 dissolved, wash the watch-glass over the dish, and evaporate 
 [Hood] the contents of the latter on the water bath, or on a 
 beaker of boiling water (Fig. 4). (Avoid loss of time by bor- 
 rowing a second Bunsen burner and gas tubing temporarily 
 from the storeroom.) Allow the dishes to cool, and weigh. To 
 make sure that the drying was complete, heat once more for 
 
9] CHEMICAL PHENOMENA 13 
 
 half an hour and weigh again. This precaution is taken in all 
 experiments of this kind. The product is common salt. Cal- 
 culate the proportion of carbonate taken to salt produced thus: 
 
 Wt. of carbonate : Wt. of salt : : 1 :x. 
 
 Compare the two values of x found from the two differing quan* 
 tities of the carbonate and interpret. 
 
 What physical property of hydrochloric acid permits use of 
 an excess of the acid without damage to the result? What 
 form of statement of the law is verified directly by this experi- 
 ment? 
 
CHAPTER HI. 
 
 OXYGEN. 
 
 10. Sources (p. 11, footnote). Heat small quantities of 
 barium peroxide, lead dioxide, potassium nftrate, silicon dioxide 
 (sand), and manganese dioxide (dry this before use by heating 
 it for a few minutes in a porcelain crucible) separately in a hard 
 glass test-tube [Note 27]. Observe whether any gas is given 
 off, and apply the test of the glowing splinter of wood [Note 28]. 
 If the Bunsen flame proves inadequate, try the blast-lamp. 
 Note any changes in appearance during the heating and de- 
 scribe the residues [Note 29]. Clean the test-tube carefully 
 with hot nitric acid [CARE] and dry it [Note 26, p. 10] after 
 each experiment. 
 
 Note 27. Use the clamp and ring-stand to support the tube, 
 or grasp it by means of a strip of folded filter paper 
 f=t (Fig. 5). In either case it must be kept in a horizontal 
 
 gj} position, etherwise condensed moisture may run down 
 and cause it to crack. 
 
 Note 28. Use a long splinter of pine wood [Side- 
 shelf] bearing a spark at one end (why not an ordinary 
 match?). Such a splinter is at once kindled by oxygen. 
 Note 29. Beginning with this chapter include in 
 p. your notes equations for all the chemical changes 
 
 you observe. When no change is observed, do not 
 attempt to give an equation until otherwise instructed 
 (Chap. IX). In the present experiments the formulae of the 
 materials used will have to be sought in the text-book. The 
 formulae of the products will also be sought in the book after the 
 physical properties of the product and the evolution or non- 
 evolution of oxygen have been noted and an indication of what 
 to seek for has thus been obtained. For other information in 
 regard to the note-book, see the "General Instructions," Notes 
 1-5, p. 1. 
 
 11. Catalytic Action. 
 
 a. Place in a test-tube a little potassium chlorate, fix the 
 tube in a vertical position by means of a burette clamp attached 
 to the ring-stand, and melt the substance at the lowest tem- 
 perature at which this is possible (patience!). Note whether 
 there is evidence of the evolution of oxygen. Now throw into 
 the molten material a pinch of powdered manganese dioxide (?) 
 [Note 2, p. 1], keeping the face out of the path of anything 
 
 14 
 
12] 
 
 OXYGEN 
 
 15 
 
 which may accidentally be projected from the tube. Interpret 
 [R 65, 75]. 
 
 6. Devise a way of showing that the manganese dioxide, 
 used as in a, remains unchanged after the action, and that it 
 is the potassium chlorate that loses its oxygen, and try it (?). 
 Consult the instructor in regard to the details of your plan 
 before executing it. 
 
 12. Preparation. Mix on paper about 5 g. of potassium 
 chlorate and 3 g. of powdered, dried (see 10) manganese dioxide. 
 Place the mixture in a hard glass test-tube provided with a one- 
 
 Fig. 6 
 
 hole cork and delivery tube (Fig. 6). Test the apparatus to 
 see that it is air-tight. (In this instance, place the end of the 
 delivery tube in the mouth, withdraw some of the air by suc- 
 tion, and note whether or not the tongue will adhere to the 
 end of the tube. If it will not, there is a leakage which must 
 be remedied.) 
 
 Clamp the test-tube on the ring-stand and heat carefully, 
 so as not to cause too violent an evolution of oxygen. Collect 
 four bottles of the gas over water in a pneumatic trough, taking 
 care to remove the delivery tube from the water as soon as the 
 bottles are full of oxygen (why?). The test-tube may be 
 cleaned by allowing it to soak in water. 
 
16 
 
 OXYGEN 
 
 [13 
 
 13. Properties [Note 5, p. 1]. 
 
 a. Place a very small amount of sulphur in a deflagrating 
 spoon, ignite it and then lower into a bottle of oxygen (?) 
 [Note 2, p. 1]. 
 
 Remove the spoon, add a little water to the jar, close the 
 mouth with the hand and shake (?). Test the water with blue 
 litmus paper or solution [R 70-71, 355]. 
 
 6. Place a little red phosphorus in the spoon, ignite it, and 
 lower into the second bottle (?). 
 
 Proceed as in a. Test with blue litmus [R 70-71]. 
 
 (If yellow phosphorus is used, it must always be cut under 
 water and handled with forceps. Great care must be taken 
 not to touch it with the hand, as it catches fire easily, and 
 causes very severe burns. Red phosphorus is safer, and should 
 be employed if available.) 
 
 c.. Lower a splinter of glowing charcoal into the third bottle, 
 holding it in the tongs or wrapping the end of a piece of copper 
 wire round it (?). 
 
 Proceed as in a (use no litmus), and cover the mouth of the 
 bottle with a glass plate. Add some lime-water (calcium 
 hydroxide) and shake again [R 481-482] (?). 
 
 14. Slow Oxidation of Metals. Devise a way of showing 
 that air loses a part of its substance (not, e.g., that the iron 
 gets heavier, but that the air diminishes in amount) when 
 moist iron powder rusts, and try it. Submit your arrangement 
 to the instructor for criticism before using it. 
 
 16. Weight of a Liter of Oxygen [Quant.] (p. 11, footnote). 
 
 Fig. 7 
 
 a. Powder some potassium chlorate; and dry it on a watch- 
 glass on the radiator, or high above a small Bunsen flame, or 
 in a drying oven. Construct an aspirator (Fig. 7), using the 
 
15 a 
 
 Wt. of tube with chlorate 
 
 Wt. of tube with residue 
 
 Wt. of oxygen 
 
 Wt. of beaker with water 
 Wt. of beaker 
 Wt. of water 
 
 Temp, in lab. 
 
 Temp, near barometer 
 
 Barometric reading 
 
 Correction 
 
 Barometer (corr.) 
 
 Aqueous tension at temp, of lab. 
 
 Partial press, ox. 
 
15] OXYGEN 17 
 
 1-liter bottle, and connect it with a hard glass test-tube. Fit 
 a nozzle to the rubber tube and slip a pinch clamp over the out- 
 let tube (syphon) in readiness for closing, at a later stage, the 
 rubber tube that connects the two glass tubes. Having inserted 
 the stopper tightly and connected the test-tube, test the appara- 
 tus to see that all the joints are air-tight [Instructions: Place 
 some water in the bottle, blow a few bubbles of air into the 
 apparatus through the syphon, and observe whether the water 
 remains permanently elevated in the vertical tube. Absolute 
 certainty that the apparatus is air-tight must be reached before 
 proceeding further]. 
 
 Carefully weigh the hard glass test-tube, then place in it 
 between 1 g. and 2 g. of potassium chlorate and weigh again. 
 More than 2 g. is not needed, less than 1 g. will not be sufficient. 
 Subtraction gives the exact weight that has been employed. 
 Fill the bottle almost, but not quite, to the shoulder with tap 
 water. After reattaching the test-tube make sure that the 
 apparatus is once more air-tight, by repeating the above- 
 described test. Next fill the syphon and nozzle completely 
 with water by blowing sharply through the latter, and close the 
 clamp. Allow the nozzle to touch the bottom of a beaker 
 (400 c.c.) containing some water. Now open the clamp and 
 raise the beaker till the levels of the water in this and the bottle 
 are the same, and the gaseous pressure therefore alike in both. 
 Close the clamp again, empty the beaker and replace it in 
 position. 
 
 Open the clamp once more and decompose the compound 
 slowly by heating, catching in the beaker the water driven 
 over by the gas. During the earlier stages a smoke, consisting 
 of solid particles, will arise. This must on no account be driven 
 into the connecting tube, as all the solid must remain in the 
 test-tube to be weighed (why?). Suspend the heating as often 
 as may be necessary to let this smoke settle. During such 
 intervals of waiting see that the nozzle is constantly immersed 
 in the water that has already passed over (why?). Stop heat- 
 ing if the tube shows signs of softening [Note 30, below] or 
 when the decomposition is complete. For the purpose of this 
 experiment (see 15 6) it is not necessary that the action should 
 be carried to completion. If, when the heating is stopped, 
 the nozzle is not under water, raise the beaker until it is well 
 covered. Allow the whole apparatus to stand until it has 
 reached the temperature of the air. Some water will return 
 to the bottle by the syphon during the cooling. Admission of 
 air to the syphon through the nozzle at any stage will prevent 
 this transference, which is essential to the success of the ex- 
 
18 OXYGEN [ 15 
 
 periment. Equalize the levels of the water in both vessels by 
 raising or lowering the beaker, and then close the clamp. 
 
 Measure the volume of water in the beaker by weighing the 
 vessel, first with and then without the water, upon the labo- 
 ratory scales. The difference in weight in grams represents 
 with sufficient accuracy the number of c.c. of water displaced 
 and hence of oxygen evolved in the operation (what of the air 
 originally in the apparatus?). Weigh the test-tube once more 
 with care [Balance]. Observe at the same time the tempera- 
 ture in the laboratory to learn the temperature of the oxygen, 
 and read the barometer [Note 31, below] to learn the pressure 
 of the air and therefore of the oxygen. 
 
 Subtract the aqueous tension (Appendix II) at the observed 
 laboratory temperature from the barometric reading (corr.) 
 to get the true (partial) pressure of oxygen in the bottle. Reduce 
 the volume by rule to and 760 mm. The weight of this 
 volume of oxygen is obtained by subtracting the weight of the 
 residue in the test-tube from the weight of the potassium 
 chlorate originally taken. Calculate by proportion from the 
 data obtained the weight of 1 liter of oxygen (x) : 
 
 Vol. of ox. found : Wt. of ox. found : : 1000 c.c. : x. 
 
 Calculate also the volume occupied by 32 g. of oxygen. To 
 what class of gases would the use of the aspirator be confined 
 for purposes like the above? 
 
 Note 30. To avoid softening of the glass, through overheating, 
 watch the color of the Bunsen flame. The blue flame is tinged 
 with a yellow color (caused by compounds of sodium in the glass) 
 where trie flame encounters the overheated part of the tube. 
 
 Note 31. When the barometer is read, the height must be 
 "corrected" to that of a column of mercury at (Appendix I). 
 
 6. Detach the hard glass test-tube from 15 a and drive off 
 the last traces of oxygen by heating strongly every portion of 
 the tube to which any of the residue adheres. Allow it to cool 
 and weigh. Obtain the weight both of the residue (potassium 
 chloride) and of the total oxygen, by difference. 
 
 Assuming the formula of the chloride to be KC1, that of the 
 chlorate is KClOar. The formula weight of KC1 being 39.15 + 
 35.45 = 74.6, find the formula weight of O* by the proportion: 
 
 Wt. of KC1 found : Wt. of oxygen found : : 74.6 : y, 
 
 where y = Oa?. This will be a multiple of 16 by a whole 
 number (why?) if the measurement has been carried out suc- 
 cessfully. What formula do you find for potassium chlorate? 
 Write the equation for the decomposition of the chlorate. 
 
156 
 
 Wt. of tube with chlorate 
 Wt. of tube empty 
 
 Wt. of chlorate 
 
 Wt. of tube with chloride 
 Wt. of tube empty. 
 
 Wt. of chloride 
 
CHAPTER IV. 
 
 HYDROGEN. 
 
 16. Interaction of Metals and Acids. 
 
 a. Place a few small pieces of each of the metals, tin (gran- 
 ulated), copper (a nail), iron (filings), zinc (gran.), lead (gran.), 
 aluminium (turnings), and magnesium (wire) in separate test- 
 tubes. Place in the measuring cylinder 20 c.c. of pure [Side- 
 shelf], concentrated hydrochloric acid [R 93] and add an equal 
 volume of water. Add 5 c.c. of the mixture to the contents of 
 each test-tube (?). Observe each case critically [Note 32, 
 below] and record the results [Note 29, p. 14], noting the order 
 of the metals in respect to activity. Notice the effect of heating, 
 if no action occurs in the cold. If heating seems to produce 
 gas, remember that it may be hydrogen chloride (why?) or 
 steam and not hydrogen. The presence of hydrogen may be 
 inferred from continued effervescence when heat is not being 
 furnished, and may be proved by the slight explosion which 
 follows when a light is brought to the mouth of the tube. But 
 hydrogen will not burn when mixed with much air or with other 
 vaporous substances (why ?). After the action has ceased, 
 filter any one of the solutions and evaporate [Hood] it to dry- 
 ness on the sand bath (?). 
 
 Can you give any grounds for the belief that the hydrogen 
 comes from the acid and not from the metal or the water? 
 
 Which metals, not used above, will displace hydrogen from 
 dilute acids [R 362, and Appendix VII] ? In what ways, if at 
 all, does the order of activity you have observed differ from 
 that accepted by chemists ? 
 
 In making the equations for these actions all that is needed, 
 beyond the information given above and acquired by obser- 
 vation, is to find the formulae of the products. These are not 
 here to be sought by measurement, but in the text-book [R]. 
 
 b. Ascertain the influence of the physical state of the metal 
 on its apparent activity by adding some zinc dust to the same 
 acid [R 111]. 
 
 c. Place one small piece of pure zinc into each of two test- 
 tubes and add diluted sulphuric acid [Desk] to both tubes. 
 If little or no action takes place in the cold, try upon one tube 
 the effect of heating (?). Try the effect of putting a platinum 
 wire in contact with the zinc in the second tube (?). Notice 
 
 19 
 
20 
 
 HYDROGEN 
 
 [17 
 
 where the hydrogen appears to come from, and explain [R 96]. 
 Withdraw the platinum wire, add a drop of cupric sulphate 
 solution, and shake (?). Is the solution still blue, and is the 
 zinc still silvery? What is the substance upon the zinc [R 361], 
 and what effect does it produce upon the apparent activity of 
 this metal ? Explain. 
 
 d. Now compare (?) the behavior of concentrated sulphuric 
 acid [Desk] with that of the diluted acid used in 16 c, by placing 
 some zinc (gran.) in a test-tube and adding enough of the 
 concentrated acid to cover the metal (why not more?). After 
 noticing the effect in the cold (?), apply heat (?). What prod- 
 ucts are formed [Note 32] ? 
 
 e. Try the interaction of acetic acid with zinc (gran.) or iron 
 (filings). Apply heat, if necessary. 
 
 Note 32. In observing an interaction a chemist first mixes the 
 substances thoroughly by shaking. If nothing occurs, he then 
 heats. If his eye detects evidence of the production of a gas or 
 vapor, he finally smells the contents of the tube. Apply these 
 three methods of observation to d before drawing any conclusion. 
 
 17. Other Methods of Obtaining Hydrogen. 
 
 a. Describe the result of throwing pieces of sodium and 
 of potassium into water, as you recall having seen it [Class- 
 room]. What other metals displace hydrogen from water 
 [R362]? 
 
 b. Fit a test-tube with a one-hole cork and delivery tube 
 (Fig. 6). Pulverize about 2 g. of sodium hydroxide, mix it 
 intimately in the mortar with about 3 g. of zinc dust, and place 
 the mixture in the test-tube. Insert the cork and delivery 
 tube, test the apparatus for air-tightness (12), and clamp the 
 tube in a horizontal position (why?). Heat the mixture and 
 
 collect the gas (?) over water (Fig. 6) 
 in a test-tube. If the tube should 
 crack [CAUTION!] cease heating at 
 once. To learn whether the gas is 
 combustible, carry the test-tube, 
 when full of the gas, mouth down- 
 ward to a flame (?). 
 
 18. Preparation of Hydrogen 
 [Same apparatus is used for 19 
 and 20]. Fit a 250 c.c. flask 
 with a safety tube and exit tube 
 (Fig. 8), and test for air-tight- 
 ness. Place in the flask some 
 
 Fig. 8 
 
 commercial, granulated zinc, close the apparatus, and attach 
 to the b-shaped exit tube, by means of a short piece of rubber 
 
19] HYDROGEN 21 
 
 tubing, a longer glass delivery tube. Now pour 30-40 c.c, of 
 dilute hydrochloric acid [Desk] through the safety tube. Test 
 the issuing gas until it is found free from air [Instructions: 
 The gas must not be ignited at once or the apparatus will be 
 blown up (whence comes the air mixed with the hydrogen?). 
 Collect samples of the gas from time to time by raising the 
 delivery tube and inserting it to the upper part (why?) of a 
 test-tube held in an inverted position, and then bringing the 
 mouth of the tube (still inverted) to a distant flame. When 
 the gas contained in the test-tube, after the first explosion of 
 the part nearest to the mouth, burns quietly, the gas is free from 
 air. Call an instructor to inspect the test]. Keep the appara- 
 tus in operation for use in 19 and 20. 
 19. Properties of Hydrogen. 
 
 a. By collection over water in the pneumatic trough, fill a 
 test-tube with hydrogen. Using this, and a similar test-tube 
 filled with air, show by an experiment that hydrogen is lighter 
 than air (?). 
 
 Fill another test-tube with hydrogen and apply it closely, 
 mouth downward, to a test-tube of air, mouth upward. Allow 
 the tubes to remain in this position for three minutes, then 
 bring first the lower and then the upper tube quickly to a 
 flame (?). What fact about diffusion does the result illustrate? 
 
 b. Fill another test-tube with hydrogen, as in a, and, hold- 
 ing it mouth downward, insert into the tube a burning match. 
 Does the match continue to burn? Explain. 
 
 c. Remove the glass delivery tube, and connect a glass 
 nozzle with the exit tube of the generating flask. Press the 
 mouth of the nozzle close against the side of a cold, dry beaker. 
 If moisture is deposited (what is its source?), fill a U-tube with 
 calcium chloride [R 100] and connect it between the exit tube 
 and the nozzle. Does the gas now deposit moisture upon a 
 beaker? If not, ascertain that the issuing gas is not explosive 
 (see test in 18), and set fire to it. What is the color of the 
 flame? Does the color change, and, if so, why? Hold a cold, 
 dry, inverted beaker over the flame. What is deposited on the 
 beaker? Why did we use dried hydrogen for this experiment? 
 Keep the apparatus in operation for 20, unless a Kipp's appa- 
 ratus is available or hydrogen is furnished in the laboratory. 
 
 Remove the nozzle with its rubber connection and push the 
 latter tightly into the end of the gas tubing furnished with the 
 Bunsen burner. Connect the other end of the tubing with 
 the illuminating-gas supply. Turn on a gentle stream of gas and 
 set fire to it. Hold a cold, dry beaker over this flame (?). 
 What inference in regard to illuininating-gas do you draw? 
 
22 HYDROGEN [ 20 
 
 J 
 
 20. Reduction by Means of Hydrogen [Two students work- 
 ing together]. Fit a hard glass tube, 25-30 cm. long and open 
 at both ends, with perforated corks and short glass tubes (Fig. 
 9). Support it on the ring-stand by means of a clamp grasping 
 
 it close to one end. Dry 1-2 g. 
 of ferric oxide by heating in a 
 porcelain crucible. Put the oxide 
 into the porcelain boat and place 
 -..-.- .M *'!=* the latter in the hard glass tube 
 near the end remote from the 
 clamp (why?), but not so near 
 that the subsequent heating will 
 char the cork. Insert the corks 
 with their tubes. Connect the 
 
 : . tube nearest to the boat with 
 
 I u \ a Kipp's apparatus and drying 
 
 / \ bottle (Fig. 18) delivering dry 
 
 Fig. p hydrogen, or with the labora- 
 
 tory supply of the gas, or with 
 the apparatus furnishing dry hydrogen used in 18. 
 
 Test the issuing gas to see that it is not explosive (see 18) 
 and repeat the test every time the apparatus is opened. Now 
 heat the boat and contents, at first gently by waving the flame 
 under the tube, and later strongly until the material is red-hot 
 (what temperature is used [R 73] ?). Observe the effect upon 
 the ferric oxide [R 756] (?). Does anything condense in the 
 cooler end of the tube? 
 
 What is the action of steam upon heated iron [R 98] ? How 
 do you reconcile this fact with that observed above? 
 
 Repeat the above experiment, using freshly heated alumin- 
 ium oxide in place of, ferric oxide, and otherwise following the 
 directions exactly (?). If no effect is noticed after an oxide 
 has been at a red heat for three minutes, absence of action 
 may be inferred. How do you account for the result obtained 
 with aluminium oxide? Which oxides are reducible by hydro- 
 gen [R 362] ? 
 
 WTiat are the commercial and mineralogical names given to 
 the two oxides used in these experiments [R] ? 
 
CHAPTER V. 
 
 WATER AND SOLUTION. 
 
 21. Purity of Water. Place a few drops of distilled water 
 on a clean watch-glass (not an evaporating-dish. Why?) and 
 evaporate on the water bath. Do the same with ordinary water. 
 Observe whether any stains remain on the glasses (?). What 
 class of impurities would leave no trace of their presence in 
 this test? 
 
 22. Union with Oxides. Place a pinch of cupric oxide in a 
 test-tube and wash it by shaking with a little distilled water 
 and pouring off the liquid. Add more water and shake again. 
 Test this latter solution with litmus paper (?). At the same 
 time test a sample of the water with litmus paper and 
 compare the tints. Repeat with barium oxide (?). 
 
 Recall the behavior of acid-forming oxides examined in 13. 
 Some oxides do not interact readily with water (which?). These 
 oxides which do interact may be divided, according to the 
 natures of the products they give, into two classes. What are 
 those classes, and which oxides belong to each? What are the 
 two classes of elements whose oxides belong to the two groups, 
 respectively [R 119]? 
 
 23. Hydrates. 
 
 a. Heat some blue vitriol (p. 4, footnote) gently in a 
 porcelain crucible (?). Allow a very small portion of the 
 white powder to stand exposed to the air on a watch-glass (?). 
 Dissolve the remainder by boiling with the minimum amount of 
 water required to dissolve it, and set the solution aside (?). What 
 chemical actions have taken place in these three operations ? 
 
 b. [Quant.] Place small quantities (about 1 g.) of Glauber's 
 salt and of blue vitriol in two porcelain dishes and ascertain 
 the gross weight in each case. Allow the dishes and contents 
 to remain for 24 hours or more and weigh again (?). Interpret 
 the results [R 121]. 
 
 c. Gently warm small quantities of barium chloride, potas- 
 sium nitrate, magnesium sulphate, and potassium dichromate 
 separately in dry test-tubes and notice whether they undergo 
 any change [R]. 
 
 Are all crystalline substances hydrates? Classify the sub- 
 stances you have examined into two groups with reference to 
 this property. 
 
 23 
 
24 WATER AND SOLUTION [ 24 
 
 d. Take in a test-tube about 5 c.c. of commercial, concen- 
 trated sulphuric acid, place in it a crystal of blue vitriol and 
 let the materials stand for an hour, or more (?) [R 388]. Now 
 heat the contents of the tube to the boiling-point of the acid, 
 holding the test-tube in a test-tube holder, keeping it far from 
 the clothing, and taking care that none of the contents spirt 
 out upon the hands or face [CAUTION: Sulphuric acid burns 
 are very painful. Note 16, p. 2]. After the contents of the 
 tube have settled, pour off the clear liquid into another tube. 
 On the following day, examine the little, shining particles on 
 the sides of the tube. What is their color, and condition? Of 
 what are they composed [R 625] ? 
 
 Are the compounds (such as cupric sulphate) which, in pres- 
 ence of water, yield crystalline hydrates, amorphous or crys- 
 talline in its absence? What is the true significance of the 
 crystalline condition [R 123] ? 
 
 e. Take a clean match-stick and, after dipping it in a solu- 
 tion of cobalt chloride, write upon a piece of white paper. After 
 the writing is dry, warm the paper gently by waving it above 
 a Bunsen flame (?). Now, breathe repeatedly upon the writ- 
 ing (?). Write equations for the actions that have occurred 
 [R 759]. 
 
 /. [Quant.] Put about 1 g. of crystals of gypsum in a weighed 
 porcelain crucible and weigh again. Place the crucible on the 
 clay triangle, heat to redness until no further loss in weight 
 occurs, and determine by difference the loss in weight (water) 
 and the weight of the calcium sulphate remaining. 
 
 The formula of gypsum must be CaSO 4 ,:rH 2 O. Assuming 
 the formula- weights of calcium sulphate (CaS0 4 = 136) and 
 of water (H 2 O = 18) calculate from your data the value of x. 
 
 Wt. of calcium sulphate: Wt. of water : : 136 : xx 18. 
 
 What is the formula of gypsum? 
 
 24. Solution: Gases in 
 Liquids. Half fill a 1-liter 
 bottle with distilled water, 
 cork, and shake vigorously 
 till the water is saturated 
 with air. Take the tempera- 
 ture of the water and also 
 the barometric reading. Fit 
 F >g- I0 a small flask (100 c.c.) with 
 
 a one-hole cork and delivery tube (Fig. 10) and measure its 
 content up to the lower surface of the cork. Completely fill 
 the whole apparatus, including the delivery tube, with the 
 
26] WATER AND SOLUTION 25 
 
 prepared water, and boil, collecting the gas in a small test- 
 tube inverted over water. When no more gas comes over, 
 equalize the levels of the water in the tube and trough (or 
 beaker) and mark the level in the tube with a thin rubber ring 
 (cut this from a piece of rubber tubing). Measure the volume 
 which the air occupied. Correct the volume for the aqueous 
 vapor present only, obtaining thus the volume of the air when 
 dry and at the observed temperature and pressure. Calculate 
 the volume of air dissolved by 100 c.c. of water at the observed 
 temperature and pressure (?). What proportion of its own 
 volume of air has the water dissolved? 
 26. Solution: Liquids in Liquids. 
 
 a. Take 5 c.c. of carbon disulphide in a dry test-tube [Note 
 26, p. 10], and add to it 5 c.c. of water, a drop at a time, shak- 
 ing vigorously after each addition and observing whether at each 
 stage the mixture is homogeneous or not (?). 
 
 b. Repeat, using 5 c.c. of ether with water (?). 
 
 c. Mark off on a narrow test-tube by means of the trian- 
 gular file the points at which 10 c.c. and 20 c.c. of water stand, 
 respectively. Empty the test-tube and pour in 10 c.c. of 
 water, and then take 10 c.c. of alcohol in the graduated cylinder 
 and add it to the water, a drop or two at a time, shaking vig- 
 orously after each addition and observing as before. Finally, 
 compare the volume of the mixture with that of the components 
 separately (?). 
 
 26. Solution: Solids in Liquids. 
 
 a. Select two large crystals of some soluble, colored com- 
 pound, such as blue vitriol (or potassium dichromate). Choose 
 crystals of approximately equal size and so large as to be just 
 capable of being slipped into a test-tube. Reduce one of the 
 crystals to a fine, uniform powder in the mortar. Note and 
 account for the change in tint (?). Now place the crystal and 
 the powder in separate test-tubes, and provide two corks which 
 fit the mouths of the tubes. Add 20 c.c. of water simulta- 
 neously to each tube, and insert the corks. Read the time 
 on a watch and shake the tubes simultaneously with vigor, 
 noting the time at which the solid finally disappears in each 
 tube (?). Account for the difference. 
 
 b. Reduce to a fine powder 10-15 g. of potassium dichromate, 
 using the larger amount in warm weather, the smaller in cool. 
 Prepare a saturated solution of the substance by placing it in 
 a flask with 50 c.c. of water and shaking at intervals for ten 
 minutes. So much of the solid must be taken that an undis- 
 solved residue remains. 
 
26 WATER AND SOLUTION [26 
 
 Take the final temperature of the solution. Now pour the 
 clear solution into a burette attached to the ring-stand, filling 
 the apparatus completely to the point of the nozzle with the 
 liquid (7 6). Read the level of the lower side of the meniscus. 
 Weigh [Quant.] a clean, dry evaporating-dish, run into it about 
 22-24 c.c. of the solution, and weigh [Quant.] the dish and con- 
 tents. Read also the level of the meniscus and note the volume 
 of solution used. Now evaporate the weighed portion of the 
 solution completely to dryness upon a water bath, or on a 
 beaker of boiling water, and weigh again. Determine by dif- 
 ference the weights of the dry dichromate here found, and of the 
 water in which it was dissolved. Calculate from the data the 
 weight of dichromate which would be dissolved by 100 c.c. of 
 water at the observed temperature. 
 
 From the volume of the part of the solution evaporated, and 
 the weight of dichromate found in it, calculate also the molar 
 solubility [R 149] of potassium dichromate at the observed 
 temperature. Compare the results with those which would be 
 obtained with potassium chromate K 2 Cr0 4 (Appendix IV) (?). 
 
 c. Take about 6 g. of the dichromate and boil with 10 c.c. 
 of water in a test-tube. Is the solubility at this temperature 
 different? Allow the clear solution to cool (?). Explain the 
 result. What sort of curve of solubility would this substance 
 exhibit (Appendix V) ? 
 
 Take about 6 g. of sodium chloride and boil with 10 c.c. of 
 water in a test-tube. Pour the clear liquid immediately into 
 another test-tube. Examine this when cool (?). Is salt much 
 less soluble in cold than in boiling water? How would its curve 
 of solubility differ from that of potassium dichromate? 
 
 d. Shake some powdered calcium sulphate with cold, recently 
 boiled, distilled water. Ascertain whether any of the salt has 
 gone into solution (21). Repeat with chalk (calcium carbo- 
 nate), rejecting the water with which it is first shaken (?). 
 Which of these substances do you find to be more soluble? 
 What are the amounts of these salts dissolved by 100 c.c. of 
 water at 18 [R Appendix IV] ? In what ratio is calcium 
 sulphate more soluble than calcium carbonate [R Appendix IV]? 
 In what ratio is calcium chloride more soluble than calcium 
 sulphate [R] ? Which of all these substances are spoken of as 
 "insoluble"? 
 
 e. Take about 10 c.c. of water in a test-tube, add to it not 
 more than 1 c.c. of lead nitrate solution, and mix. Now add 
 about 2 c.c. of dilute hydrochloric acid (?). Repeat, heating 
 the mixture to the boiling-point before adding the hydrochloric 
 acid (?). Examine this tube again, after the contents have 
 
27] WATER AND SOLUTION 27 
 
 cooled (?). Interpret the result. Is lead chloride an "insol- 
 uble" substance [R Appendix IV]? 
 
 /. Take about 10 c.c. of water in each of two test-tubes. 
 To one portion add Glauber's salt, previously pulverized in a 
 mortar, until, after shaking, a considerable excess remains un- 
 dissolved. Saturate the other portion with anhydrous sodium 
 sulphate in the same manner. Perform the last operation 
 rapidly, taking care not to introduce any particles of the hydrate, 
 and do not let the solution stand before use. Now decant the 
 two liquids into clean test-tubes, disregarding the cloudiness 
 of one of them. Then add a little of the anhydrous substance 
 to the solution first made and a small crystal of Glauber's 
 salt to the contents of the second test-tube and shake both (?). 
 After a short time, examine the contents of each again (?). 
 Interpret the results [R 160-161]. 
 
 g. Two immiscible solvents. Place one small particle of 
 iodine in each of three test-tubes and add to one water, to the 
 second potassium iodide solution, to the third carbon disul- 
 phide, and shake each (?). If any iodine remains undissolved, 
 pour off that solution into a clean test-tube. Now add a drop 
 or two of carbon disulphide to the first two solutions, shake again 
 (?), and describe carefully what seems to have happened. Deduce 
 from this the relative solubility of iodine in the three solvents. 
 
 27. Properties of Solutions: Vapor pressure and Boiling- 
 Point. 
 
 a. Place some dry potassium carbonate (or pulverized cal- 
 cium chloride) in a small beaker or crucible. Set the vessel in 
 an evaporating dish containing water, and invert over it a 
 larger beaker so that the edge of the latter is under the liquid. 
 Examine the material from day to day (?). Remembering that 
 there is moisture upon the surface of even "dry" bodies, and 
 that therefore a solution of potassium carbonate was present 
 with the solid, explain the change [R 162]. To what class of 
 substances do those which deliquesce all belong? Is deliques- 
 cence a physical or a chemical phenomenon? 
 
 b. Fix a test-tube containing about 10 c.c. of water hi a 
 clamp upon the ring-stand. Suspend the thermometer from a 
 ring by means of a thread, hi such a way that the bulb is im- 
 mersed in the water. Boil the water, using a small Bunsen flame, 
 and read the temperature (?). Now add to the boiling water 
 about 2-3 g. of dry calcium chloride, and, after solution is com- 
 plete, read the temperature of boiling again (?). Add another, 
 equal portion of calcium chloride and repeat the temperature 
 reading after the whole has dissolved (?). Explain [R 162]. 
 
 c. If pulverized ice were to be added to water until the 
 
28 WATER AND SOLUTION [ 28 
 
 solid no longer melted, what would be the temperature of the 
 mixture? If ice were to be added to the aqueous solution of 
 some substance, until the ice no longer melted, how would the 
 temperature differ from that of water and ice? Why does this 
 difference exist [R 163] ? Why does salt thrown upon ice cause 
 the latter to melt [R 164] ? 
 
 28. Properties of Solutions: Volume Changes and Thermal 
 Effects. 
 
 a. Recall the change in volume observed in 25 c when alcohol 
 and water were mixed (?). 
 
 b. [Quant.]. Take about 25 g. of potassium carbonate and 
 determine its weight to the nearest tenth of a gram. Assum- 
 ing the specific gravity of this substance to be 2, calculate the 
 volume of the amount you have taken (?). Place in the 
 graduated cylinder exactly 85 c.c. of water and take its tem- 
 perature. What is the sum of the volumes of the water and the 
 carbonate, separately? Add the weighed specimen of potas- 
 sium carbonate to the water, dissolve by repeated inversion 
 of the cylinder, closing the mouth of the latter with the hand, 
 and read the volume of the solution (?). Read also the tem- 
 perature of the solution immediately. Is there a change in 
 volume, or in temperature, on dissolving two substances in one 
 another? 
 
 What relation exists between the sign of the thermal effect 
 when a substance is dissolved in a nearly saturated solution of 
 the same substance, and the change of solubility with tem- 
 perature [R 260] ? What do you infer in this case? 
 
 c. Examine the solubility curve of anhydrous sodium sul- 
 phate [R 158] (?). Will this compound give out or absorb 
 heat in dissolving in water [R 260] ? Verify your conclusion by 
 trying the experiment (?). 
 
 d. Repeat b, using about 25 g. of ammonium chloride (sp. 
 gr. 1.5). Make the same observations and answer the same 
 questions. 
 
CHAPTER VI. 
 
 CHLORINE AND HYDROGEN CHLORIDE. 
 
 29. Preparation of Chlorine [Hood]. 
 
 Experiments 29 b and 30 must be accomplished at ojje ex- 
 ercise. 29 a may be postponed to facilitate this. 
 
 a. Prepare some strips of filter paper by dipping them in 
 starch emulsion [Side-shelf] to ^ich you have added one drop 
 of potassium iodide solution. 
 
 Place small quantities of finely powdered manganese dioxide, 
 potassium chlorate, lead dioxide, and pure litharge in as many 
 test-tubes, and add a little commercial, concentrated hydro- 
 chloric acid [Desk] to each. Notice the color (?) and odor (?) 
 of the gas in each case. If no action takes place in the cold, 
 apply heat. Dip into the gas in one of the test-tubes a strip 
 of the prepared paper (?). How do you account for the differ- 
 ence in the behavior of the two oxides of lead? Do all com- 
 pounds containing oxygen give free chlorine in this way? If 
 not, state what is common to those which do. 
 
 6. Fit up a 250 c.c. generating flask, as in Fig. 11 with a 
 dropping-funnel (or substitute 3(5 fr) 
 and an L-shaped glass tube at- 
 tached to the exit tube. Use the 
 shortest possible rubber connec- 
 tions here and in 31 (hydrogen 
 chloride), as rubber tubing is 
 destroyed by these gases. Test 
 the apparatus to see that it is air- 
 tight. Place in the flask about 
 20 g. of dry potassium permanga- 
 nate, and fill the globe of the 
 dropping-funnel with diluted, 
 commercial, concentrated hydro- 
 chloric acid (1 Aq. : 3 acid). Allow the delivery tube to dip 
 to the bottom of a small beaker containing a little sodium 
 hydroxide solution [Desk]. Now admit the acid drop by drop, 
 regulating the flow so that too rapid a stream of gas is not 
 produced. The complete displacement of the air in the flask 
 will be recognized by the color of the contents and the fact that 
 the bubbles of pure chlorine are completely absorbed by the 
 sodium hydroxide. 
 
 29 
 
 Fig. ii 
 
30 CHLORINE [ 31 
 
 When the air hag all been displaced, fill three dry bottles and 
 one dry test-tube with the gas by downward displacement, 
 observing the following precautions: Provide a piece of stiff 
 paper or card, perforated with a hole for the reception of the 
 delivery tube, to cover the bottles during the filling, and cover 
 the vessels with glass plates as soon as they are full. See that 
 the delivery tube reaches to the bottom (why?) of each vessel 
 during the filling. Replace the end of the delivery tube in the 
 sodium hydroxide when all the vessels have been filled. When 
 the experiment is over, pour the contents of the generating 
 flask into the sink in the hood, and not into one of the sinks in 
 the open laboratory, and wash down with much water. 
 
 30. Properties of Chlorine [Hood]. 
 
 o. In one bottle of the gas scatter a pinch of finely pow- 
 dered antimony [R 715] (?). 
 
 6. Take a clean piece of sodium [From Instructor] and cut 
 from it a very thin slice not more than one-half inch square 
 (fingers and knife used in handling sodium must be dry!). In- 
 troduce this piece into a bottle of chlorine (?) and cover at once 
 with a glass plate. Examine after half an hour. If any of the 
 metal remains unattacked, scrape off the white deposit and 
 place it upon a watch-glass, and throw the sodium into the 
 sink in the hood. Add the material on the watch-glass to 
 that in the bottle and dissolve in 2 c.c. of water. Allow the 
 solution to stand in a watch-glass until it dries, and examine 
 the crystals with a lens (?). 
 
 c. Connect a glass nozzle with the illuminating-gas supply, 
 and lower a small, burning gas-jet into the third bottle (?). 
 Blow the breath into the bottle after withdrawing the jet (?). 
 
 d. Fill a test-tube with hydrogen from a Kipp's apparatus 
 or from the laboratory supply. Bring this tube mouth to 
 mouth with a tube of chlorine and mix the gases by repeated 
 inversion. (Take care not to expose the mixture to direct sun- 
 light.) Hold the mouth of each tube to the Bunsen flame (?). 
 Close the mouth of one tube quickly with the thumb, add a 
 few drops of water, shake, and test the solution with litmus 
 paper (?). 
 
 31. Preparation of Hydrogen Chloride. 
 
 a. [Hood] Place small quantities of ammonium chloride, 
 barium chloride, mercuric chloride, and sodium chloride in as 
 many test-tubes, and add a few drops of concentrated sulphuric 
 acid [Desk] to each (?). Describe what happens in each case. 
 Blow moist air across the mouth of the test-tube (?). Lower a 
 glass rod dipped in ammonium hydroxide solution into each 
 [Note 33, below]. Try the effect of heating. Remember that 
 
32] CHLORINE 31 
 
 the solubility (physical) of the substance in sulphuric acid will 
 largely determine the speed of the action. In case of difficulty, 
 therefore, take a fresh sample of the solid, pulverize it finely, 
 and shake with the acid for some minutes before heating and 
 testing. Arrange the substances in the order of apparent 
 activity. 
 
 6. To a pinch of finely powdered sodium chloride add a 
 strong solution of phosphoric acid and heat if necessary [R 179] 
 (?). Test with ammonia as before. The above remark about 
 -olubility applies also to this case. 
 
 Why is hydrogen chloride displaced completely in a and b 
 by these acids under these conditions [R 180]? 
 
 c. [Hood] In a 250 c.c. flask (Fig. 11), fitted with dropping- 
 funnel (or substitute, 36 6) and L-shaped delivery tube, place 
 about 30 g. of common salt. Admit concentrated sulphuric 
 acid through the funnel. Collect the gas in three dry bottles 
 by downward displacement (using a perforated square of paper 
 or card as in the case of chlorine), cover, when filled, with glass 
 plates and reserve for 32. Place about 10 c.c. of distilled 
 water in a test-tube, attach a nozzle to the delivery tube, and 
 allow the gas to bubble into this for a few minutes (?). Re- 
 serve the aqueous solution also for 32. Write equations for the 
 two possible interactions of salt and sulphuric acid. Which of 
 the two takes place under the above conditions? 
 
 Note 33. The use of ammonia is not a specific test for hydro- 
 gen chloride. It can be used only for ascertaining the presence or 
 absence of any one of several gases, usually of acidic character, 
 which are capable of uniting with ammonia. 
 
 In writing the equation here, and whenever the same test is 
 used, consider whether the interaction took place with the liquid 
 on the rod (containing NH 4 OH), or with the gas (NHJ given off by 
 this liquid. Note the odor of this gas (?). 
 
 32. Properties of Hydrogen Chloride and of Hydrochloric 
 Acid. 
 
 a. Invert one of the bottles of the gas in a dish of water (?). 
 Relate this property to that observed on blowing moist air 
 into the gas (31 a). If any gas remains, what should you 
 expect it to be? Test your conclusion experimentally. 
 
 6. Pour a little ammonium hydroxide solution on a strip of 
 filter paper and plunge this into the second bottle (?) [Note 
 33, above]. Describe the difference between these fumes and 
 those formed by the action of moist air upon the gas (31 a) . 
 
 c. Devise a way of proving, in a rough way, that the gas is 
 heavier than air, and use the third bottle of gas for carrying it 
 out. 
 
32 HYDROGEN CHLORIDE [ 33 
 
 Perform the following experiments with the aqueous solu- 
 tion prepared in 31 c. 
 
 d. Test the solution with litmus paper. 
 
 e. To a part add a granule of zinc (?). 
 
 /. To a part add a crystal of sodium carbonate [R 480] (?). 
 
 g. Dilute the remaining portion of the acid solution with an 
 equal volume of water and distribute it between three test- 
 tubes. To one test-tube add a drop or two of mercurous nitrate 
 solution (?), to the second a drop of lead nitrate solution (?), and 
 to the third a drop of silver nitrate solution (?). After allowing 
 the contents of each tube to settle, pour away the liquid, add 
 water, and boil (?). Allow to cool, and note, by the appearance 
 of crystals, which of the precipitates is soluble in hot water. 
 
 These precipitates are given, not only by hydrochloric acid, 
 but by any chloride, and are, therefore, means of recognizing 
 the presence of the chloride radical which is common to all 
 chlorides. 
 
 33. Theory of the Method Used in Preparing Hydrogen 
 Chloride. To a concentrated solution of sodium-hydrogen 
 sulphate add pure concentrated hydrochloric acid (?). Add 
 the acid a very little at a time to avoid over-rapid precipita- 
 tion, and agitate between additions. The longer the operation 
 takes, the better. Examine the result with a lens (?). 
 
 Write the equation for this action. What relation does this 
 action bear to that in 31 c? What circumstances determine 
 the direction of a reversible action [R 180]? 
 
CHAPTER VII. 
 
 EQUIVALENT WEIGHTS, FORMULA, EQUATIONS. 
 
 34. Composition of Carbon Dioxide [Quant. Two students 
 working together]. Fit a piece of hard glass tubing, 25-30 cm. 
 long, with perforated corks. Insert at one end a short piece 
 of glass tubing and at the other a U-tube, as in Fig. 12. The 
 
 Fig. 12 
 
 inner edges of the hard glass tube should be rounded with a 
 file, or flared by use of the blast-lamp and a piece of charcoal. 
 Rubber stoppers will give tight joints more surely than corks. 
 Attach to the U-tube by means of a cork a short, straight tube, 
 of the diameter of a narrow test-tube, which has been drawn 
 out [Bunsen flame] so as to leave a small opening at the free 
 end. Arrange a loop of thread [Side-shelf] with which to sus- 
 pend the U-tube from the balance. Place in the hard glass 
 tube a plug of granular cupric oxide, about 4 cm. in length. 
 This may be held in position by small wads of asbestos. The 
 cupric oxide and asbestos must be dried by heating in the 
 porcelain crucible before use. Put about 0.2 g. of pure dry 
 sugar-charcoal [Instructor] in a porcelain boat, weigh the 
 boat with contents, and set it in the tube close behind the cupric 
 oxide. Make a few c.c. of a strong (approximately 30 per 
 cent) solution of potassium hydroxide, fill the bend of the 
 U-tube with it, and charge the small tube beyond with frag- 
 ments of solid caustic potash. Test the apparatus with the 
 greatest care to see that it is absolutely air-tight [Instructions]. 
 Finally, immediately before starting the combustion, weigh the 
 connected potash tubes, replace them in position, and test 
 once more for air-tightness. 
 
 Slip a cylinder, made by rolling a piece of wire gauze, over 
 the hard glass tube, connect the latter with the oxygen cylin- 
 der (or other source of oxygen), and heat the part containing 
 
 33 
 
34 EQUIVALENT WEIGHTS [35 
 
 the boat and cupric oxide with two burners (extra Bunsen 
 burner from storeroom). Turn on the oxygen with care and 
 regulate the stream so that the carbon may burn slowly and 
 not more than 15-20 bubbles of unused oxygen escape per 
 minute. A more rapid stream will involve the loss of carbon 
 dioxide. Heat the front of the boat first and let the glow, 
 caused by the combustion, travel along. The burning will take 
 30-45 minutes. Continue the stream of oxygen for 4-5 minutes 
 after the carbon is completely burned (why?), then disconnect 
 the potash apparatus and weigh it. A more accurate result is 
 obtained by finally displacing the oxygen by air (why?). After 
 the tube has cooled, weigh the boat with any ash it may con- 
 tain. Return the cupric oxide to the bottle. 
 
 The loss in weight of the boat gives the amount of car- 
 bon; the gain in weight of the potash apparatus, the amount 
 of the carbon dioxide. The difference of these two gives the 
 oxygen. 
 
 Calculate from your data the weight of carbon combining 
 with 8 parts of oxygen. 
 
 Wt. of ox. found : Wt. of carbon found : : 8 : x. 
 
 This gives x, the combining weight [R 46] of carbon. This 
 weight of carbon is equivalent [R 49] to those amounts of other 
 elements which likewise unite with 8 parts of oxygen. 
 
 Assuming the atomic weights of carbon and oxygen to be 
 12 and 16 respectively, calculate from your data the formula of 
 carbon dioxide [R 57]. 
 
 Make the equation representing the action. 
 
 36. Composition of an Oxide of a Metal [Quant. Note 22, 
 p. 7]. On account of the difficulties attending the making, 
 the collecting, or the weighing of most oxides formed by direct 
 union, the following indirect method is suggested. It consists 
 in converting a known weight of a metal into the nitrate of the 
 metal by the action of nitric acid, and obtaining the oxide by 
 decomposition of this salt. Attempt no equations for these 
 actions. 
 
 a. Composition of an oxide of iron. Weigh an evaporating- 
 dish of medium size, place in it about 1 g. (12 inches) of pure 
 iron wire, and weigh again. Cover the dish with a watch-glass, 
 convex side downward, and add 10 c.c. of pure dilute nitric 
 acid. Set the dish, covered, on the water bath until the iron 
 has dissolved, adding a few drops of concentrated nitric acid 
 if any of the wire remains unattacked [Instructions]. Then 
 rinse the cover-glass carefully into the dish and remove it, and 
 evaporate the solution to dryness on the water bath or on a 
 
36] EQUIVALENT WEIGHTS 35 
 
 beaker of boiling water [Hood]. When the residue (what is 
 it?) is dry, place the dish on a clay triangle supported on the 
 ring-stand and heat carefully with a burner held in the hand as 
 long as any red fumes are given off. During this process, and 
 especially at first, great care and patience must be exercised, 
 as too rapid heating may cause solid particles of the product to 
 be thrown from the vessel. If any crackling noise is observed, 
 remove the burner at once. When, after final strong heating 
 of every part of the material, red fumes or other evidence of 
 continued change can no longer be perceived, allow the dish 
 and contents (?J to cool and weigh them. To make sure that 
 the decomposition was complete, heat once more, cool, and 
 weigh again. This precaution is always necessary in experi- 
 ments of this nature. 
 
 The difference of the weights of the metal and of the oxide 
 gives the weight of oxygen combined with the known weight 
 of iron. 
 
 Calculate from your data the weight of iron (x) combining 
 with 8 parts of oxygen. 
 
 Wt. of ox. found : Wt. of iron : : 8 : x. 
 
 This is the combining weight [R 46] of iron in this oxide. 
 This weight of iron is equivalent [R 49] to those amounts of 
 other elements which likewise unite with 8 parts of oxygen. 
 
 Assuming the atomic weights of iron and of oxygen to be 55.9 
 and 16, respectively, calculate from your data the formula of 
 the oxide of iron [R 57]. What is the name of this oxide? What 
 other oxides of iron are known [R] ? 
 
 b. Pure zinc (about 1 g.) may be used instead of iron. The 
 manipulation is the same as in 35 a. The residue from evapo- 
 ration is a syrup-like body, however, which cannot be dried. 
 Extra caution must, therefore, be used in heating this to avoid 
 loss by spirting. 
 
 Calculate the combining (equivalent) weight of zinc and the 
 formula of the oxide as in 35 a. 
 
 c. Magnesium wire (about 1 g.) may be used instead of zine 
 (35 6). Equal care is required during the evaporation and 
 heating. Tin (about 1 g.) may also be used. 
 
 36. Equivalent Weight of a Metal by Displacing Hydrogen 
 [Quant.]. 
 
 a. First fill the pneumatic trough and 1-liter bottle with 
 water so that the liquid may acquire the temperature of the 
 room. Fit a 100 c.c. flask somewhat as in Fig. 8, using in place 
 of the thistle-tube a dropping-funnel (see 36 6). To carry a 
 doubly-bored cork the mouth of the flask must be rather wide. 
 
36 EQUIVALENT WEIGHTS [36 
 
 A larger flask than 100 c.c. must not be employed on account 
 of the waste of acid its use would entail. The stem of the 
 dropping-funnel must reach to the bottom of the flask, and the 
 inner end of the L-tube must not project below the bottom of 
 the cork. Attach a rubber or glass delivery tube and see that 
 the apparatus is air-tight. Weigh a piece of chemically pure 
 zinc, -taking about 2 g. Without detaching your platinum wire 
 from the glass rod, wrap it tightly round the zinc (why)? and 
 allow the whole to slide gently into the flask. Fill the appara- 
 tus completely, from the top of the stem of the funnel to the 
 tip of the delivery tube, with water. Close the stopcock when 
 the bulb has almost emptied itself. Invert the 1-liter bottle, 
 filled with water, on the shelf of the pneumatic trough, and put 
 the delivery tube in position. 
 
 Fill the globe of the funnel with pure concentrated hydro- 
 chloric acid [Side-shelf] and admit this to the flask, a little at 
 a time, in such a way that a steady, but not too violent, action 
 takes place. A good deal will be needed at first before suffi- 
 ciently rapid action begins (why?). When the metal is entirely 
 dissolved, drive all the gas over into the bottle by pouring water 
 once more through the funnel (be careful that no air is carried 
 over with the water). 
 
 The weight of the hydrogen which has been displaced by the 
 weighed quantity of zinc is to be ascertained in this experiment, 
 not by direct weighing, but, as follows, by measuring the volume 
 of the gas and calculating its weight from this and the density 
 of hydrogen : 
 
 When the gas has acquired the temperature of the water and 
 room, lower the bottle until the level of the water outside and 
 inside is the same. If there is still a good deal of water in the 
 bottle, the latter may have to be inclined to accomplish this 
 and the following operation. Close the bottle with a cork while 
 it is in this position and remove it from the trough. To find 
 the volume of the gas, weigh (to the nearest gram) the bottle 
 as it stands on the laboratory scales, and also completely filled 
 with water, and subtract. Since 1 c.c. of water weighs approx- 
 imately 1 g. the number of grams in this difference gives at once 
 the number of c.c. of hydrogen. Why is this method of meas- 
 uring the volume of the water more accurate than using the 
 graduated cylinder? Record the temperature of the hydrogen 
 by reading a thermometer hung in the laboratory close to the 
 apparatus. Read the barometer, and record the corrected 
 reading [Note 31, p. 18]. Since the hydrogen is mixed with 
 water vapor, correct for the latter by subtracting the aqueous 
 tension (Appendix II) at the observed laboratory temperature 
 
38] EQUIVALENT WEIGHTS 37 
 
 from the corrected barometric reading. This gives the true 
 partial pressure of the hydrogen. 
 
 Reduce the volume of the hydrogen, by rule, from the ob- 
 served temperature and partial pressure to and 760 mm. 
 Find by calculation (1 liter weighs 0.09 gr. at and 760 mm.) 
 the weight of the hydrogen obtained. Calculate the equi- 
 valent weight of zinc, i.e., the weight of the metal (x) which 
 displaces 1.008 g. of hydrogen: 
 
 Wt. of hyd. : Wt. of zinc : : 1.008 : x. 
 
 Calculate from the result how many equivalent weights of 
 hydrogen would be displaced by 65.4 g. ; the gram-atomic 
 weight, of zinc. What is the valence of zinc? How many for- 
 mula-weights of hydrogen chloride are required to furnish the 
 hydrogen displaceable by one atomic weight of zinc? Write 
 the left-hand side of the equation in accordance with this 
 result. 
 
 Wash the trough carefully until it is absolutely free from 
 acid, and put it away in an inverted position to avoid rusting. 
 
 b. The above experiment may be performed with magne- 
 sium (about 0.9 g.), iron (about 1. 8g.) y or aluminium (about 0.8 
 g.) in place of zinc. 
 
 In the absence of a dropping-funnel, a substitute may be 
 made by connecting a funnel with a straight tube by means of 
 a rubber joint closed with a pinch clamp. Or, the aspirator 
 (Fig. 7) may be used here, the metal (half the above quantities), 
 water, and a smaller tube containing the acid being placed in 
 the test-tube, and the mixing being effected by inclining the 
 bottle after the apparatus is connected. 
 
 37, Inter -Equivalence of Equivalent Weights (Law of Recip- 
 rocal Proportions) . If 35 b was performed, compare the quan- 
 tity of zinc which combined with 8 g. of oxygen with that which 
 in 36 a was found to displace 1.008 g. of hydrogen. If they are 
 identical, these two quantities of oxygen and hydrogen are 
 chemically equivalent and may combine with one another. 
 What substance has precisely this composition? 
 
 If magnesium was used, in 36 b, compare the quantity with 
 that combining, in 35 c, with 8 g. of oxygen and answer the 
 same question. 
 
 38, Combining Weights of Zinc and Chlorine [Quant.]. 
 (From Torrey's Elementary Studies.) Weigh an evaporating- 
 dish of medium size and place in it about 2 g. of pure zinc. Add 
 a little diluted (1 Aq. : 2 acid), pure, concentrated hydrochloric 
 acid [Side-shelf] and cover with a watch-glass, convex side 
 downward. If the action is very slow, the tip of a platinum 
 
38 EQUIVALENT WEIGHTS [ 39 
 
 wire may be placed in contact with the zinc. Maintain a brisk 
 action by further additions of concentrated hydrochloric acid 
 in very small amounts at a time. Final excess of the acid should 
 be avoided, as time will be lost in the subsequent evaporation. 
 When the metal is completely dissolved, rinse the cover-glass 
 and platinum wire carefully into the porcelain dish and remove 
 them. Allow the solution to evaporate as far as possible on 
 the water bath or on a beaker of boiling water [Hood]. Now 
 place the dish on the ring-stand, and, using a small Bunsen 
 flame, allow the syrup-like solution to evaporate slowly to 
 dryness. Then heat the white mass to the point at which 
 it has completely melted and no further. The best way to 
 achieve this with the minimum rise in temperature is to let the 
 Bunsen flame play on the surface from above. Overheating 
 must be avoided, because the product is volatile at high tem- 
 peratures. The moment the dish has so far cooled that the 
 hand can be borne upon the bottom, wipe the dish carefully 
 and weigh it. The substance absorbs moisture greedily from 
 the atmosphere, hence expedition is required in cooling and 
 weighing if accurate results are to be obtained. To insure 
 accuracy, the melting, cooling, and weighing should be repeated, 
 and the lower result taken as correct. 
 
 Calculate, from the data obtained in this experiment, how 
 much chlorine combines with the equivalent weight of zinc 
 found in 36 a. This amount is the equivalent weight of 
 chlorine (x) : 
 
 Wt. of zinc : Wt. of chlor. : : Equiv. of zinc : x. 
 
 Assuming the atomic weights of zinc and chlorine to be 65.4 
 and 35.45 respectively, determine the formula of zinc chloride. 
 
 Express the whole action of hydrochloric acid on zinc in 
 symbols by making the equation in accordance with these 
 conclusions. Which of the factors in the equation have you 
 determined experimentally, and which not? What law do we 
 use in assuming that the undetermined factors are correct? 
 
 39. Combining Weights of Lead or Silver and Chlorine. 
 From the weights of lead (9 a) or silver (96) and chlorine 
 found to combine, calculate the formula of lead (or silver) 
 chloride. 
 
 40, Multiple Proportions [Quant.]. 
 
 a. [Two students working together]. Fit a hard glass tube 
 with corks through which pass short pieces of narrow glass 
 tubing (Fig. 9). Fix the tube in a clamp upon the ring-stand, 
 attaching the clamp close to one end so that the rubber cover- 
 ing of the clamp may be away from the central part which is to 
 
40] EQUIVALENT WEIGHTS 39 
 
 be heated. Make sure that the apparatus is air-tight. Dry 
 some pulverized cupric oxide by heating it in a hard glass test- 
 tube. Weigh two clean, dry, porcelain boats, and place in one 
 about 1.5 g. of the cupric oxide and in the other 2.5 g. or more 
 of cuprous oxide * [From Instructor] and weigh each again. In 
 recording the weights and in handling the boats be careful to 
 distinguish the one from the other. Place the boats in the 
 hard glass tube, so that the points of the boats touch in the 
 center of the tube. Connect the end nearest to the cuprous 
 oxide with a source of dry hydrogen (Kipp's apparatus and 
 drying-bottle, or laboratory supply). 
 
 Pass a gentle stream of hydrogen through the apparatus 
 until a test (?) shows that the air has all been displaced. Re- 
 duce (why?) the speed of the gas until the bubbles in the wash- 
 ing-bottle can easily be counted. Now heat the boats moder- 
 ately, beginning with that containing the cuprous oxide. What 
 collects in the cooler end of the tube? Where does it come 
 from? When the action, which requires 10-15 minutes, is 
 finished, allow the boats to cool in a stream of hydrogen. Weigh 
 the boats and contents (?), taking care not to interchange 
 them. To ascertain whether the action is complete, heat the 
 boats once more in hydrogen, cool, and weigh again. 
 
 Determine by difference the weights of oxygen and copper in 
 each case, and calculate from the data how much copper is* 
 combined with 8 parts of oxygen in each of the two oxides. 
 What is the ratio of the two values of copper? Show that the 
 result illustrates the law of multiple proportions. 
 
 Assuming 16 and 63.6 to be the atomic weights of oxygen 
 and of copper, respectively, calculate from the data the formula 
 of each oxide. Construct the equations representing the 
 action of hydrogen upon each oxide. 
 
 6. Pure lead dioxide and pure lead monoxide may be used 
 as in a, if available. The monoxide, however, absorbs carbon 
 dioxide readily from the air, and therefore does not keep well. 
 The experiment must, therefore, be tried by the instructor 
 before the material on hand is given to the class. Lead mon- 
 oxide is more difficult to reduce than the dioxide. The heat- 
 ing in a stream of hydrogen must be repeated to constant 
 weight. 
 
 c. If the quantitative decomposition of potassium chlorate 
 (15 6) was carried out, dried potassium perchlorate may be 
 decomposed in the same way and the results compared. Weigh 
 an open, long, hard glass test-tube. Place in it about 1 g. of 
 
 * Pure cuprous oxide (Kahlbaum's) can be kept successfully if sealed 
 up in small bottles which are not opened until needed. 
 
40 EQUIVALENT WEIGHTS [ 41 
 
 potassium perchlorate and weigh again. Then heat the tube 
 and drive off the oxygen slowly, and weigh again. Heat 
 strongly once more, to constant weight. Ascertain by differ- 
 ence the weights of potassium chloride (residue) and of oxygen. 
 
 The formula of the perchlorate is KC10 X /. Calculate from 
 your data how much oxygen ((V) is combined with 74.6 g., the 
 formula-weight, of potassium chloride (KC1). This result, the 
 formula-weight of the oxygen, will be a multiple of 16. Com- 
 pare this formula-weight of oxygen with that found in the 
 experiment with potassium chlorate, and show how the results 
 illustrate the law of multiple proportions. Find the formula 
 of potassium perchlorate and write the equation for the decom- 
 position of the compound. 
 
 41. Dulong and Petit's Law. According to Dulong and 
 Petit [R 211], if the correct atomic weight, when it has been 
 found, is multiplied by the specific heat of the element in the 
 solid form, the product is a number which in most cases lies 
 between 6 and 6.8. 
 
 Take the values of such combining weights or equivalents 
 as you have found experimentally, viz., iron (36 a, 36 6), zinc 
 (35 6, 36 a); magnesium (35 c, 36 6), aluminium (36 6), silver 
 (39), copper (40 a, two values), or lead (39, one value, and 40 b, 
 two values), and multiply each by the corresponding specific 
 "heat (Appendix III). If the result is about 6.4, the atomic 
 weight is the same as the equivalent weight. If not, multiply 
 the equivalent weight by the smallest integer which will bring 
 the final product within the limits 6 to 6.8. The integer used 
 is the valence of the element, and the product of the equivalent 
 weight and the valence is the atomic weight. Show the working 
 hi your notes and give a list of the atomic weights and valences 
 found. In the case of copper or lead, distinguish the valences 
 in the two compounds. 
 
CHAPTER VIII. 
 
 BROMINE, IODINE, FLUORINE, AND THEIR COMPOUNDS 
 WITH HYDROGEN. 
 
 CHLORINE (Chap. VI) and the elements to be studied in this 
 chapter form a group having very similar properties, and are 
 called the halogens. Recall the facts about chlorine and 
 hydrogen chloride, and use them as a guide in trying to under- 
 stand the chemistry of the rest of the group. Remember 
 particularly that chlorine is colored, has a powerful odor, and 
 does not cause fumes in moist air; and that hydrogen chloride 
 is colorless, and causes dense fumes in moist air. The corre- 
 sponding substances throughout the group may be expected to 
 present properties like these. Thus, the elements are all colored 
 substances, the hydrogen compounds are all colorless and fume 
 in moist air. The hydrogen compounds, hydrogen chloride, 
 hydrogen bromide, etc., are known as the hydrogen halides. 
 
 42. Preparation of Bromine [Hood], Powder about 1 g. of 
 potassium bromide, mix it in the mortar intimately with about 
 2 g. of pulverized manganese dioxide, and place the mixture in 
 a test-tube. In a second test-tube dilute 2-3 c.c. of concen- 
 trated sulphuric acid [Desk] with half its volume of water (add the 
 acid, cautiously, to the water), and mix with the contents of the 
 first test-tube enough of this solution to moisten the materials 
 thoroughly, and no more. After allowing the mass to stand 
 for a few minutes, apply a gentle heat to the tube, and note 
 the color and behavior of the vapors evolved (?). Apply a 
 strip of filter paper moistened with a starch-potassium iodide 
 emulsion (29 a) to the mouth of the tube (?). 
 
 What other materials might be substituted for the potassium 
 bromide in the above experiment? 
 
 43. Properties of Bromine. 
 
 Shake up one drop of bromine [CARE. Do not spill upon the 
 hands] with 10 to 15 c.c. of water in a test-tube, and divide the 
 solution ("bromine-water") between four test-tubes. 
 
 a. To one of these add 1-2 c.c. of ether, and shake (?). Note 
 the relative solubility of bromine in ether and in water as dis- 
 played by the depth of color in each layer (26 g}. To the 
 second add carbon disulphide (?), and to the third chloro- 
 form (?), observing as before. To about 10 c.c. of starch 
 
 41 
 
42 BROMINE, IODINE, FLUORINE [ 44 
 
 emulsion add a few drops from the fourth test-tube (?) [R 230 
 and see result of 47 b, below], 
 
 b. Fit up an apparatus to generate a small amount of chlorine 
 [Hood], using a side-neck test-tube instead of the flask in Fig. 
 11 (p. 29). If chlorine-water is available, it may be used 
 instead of the gas. 
 
 Dissolve a single, very small crystal of potassium bromide in 
 a few c.c. of water in a test-tube. Add several drops of carbon 
 disulphide, and then pass a few bubbles of chlorine through the 
 solution, or add a few c.c. of chlorine-water (?). Shake, and 
 notice the appearance of color in the carbon disulphide (?). 
 Infer from this result 'the relative activities of chlorine and 
 bromine (?). The result measures the relative affinity of 
 chlorine and bromine for what elements? 
 
 The chlorine generator, if used, is required again in 47 c, 
 which may be performed at once, before the apparatus is taken 
 apart and cleaned. Prepare also 5 c.c. of saturated chlorine- 
 water, if not furnished on the side-shelf, cork it up in a test- 
 tube, and set it aside in a dark place for use in 46 g and 50 a. 
 44. Preparation of Hydrogen Bromide. 
 
 o. Pulverize about 1 g. of potassium bromide, place it in a 
 test-tube, and cover with concentrated phosphoric acid solu- 
 tion. Notice the apparent slowness of the action on account 
 of the insolubility (physical) of the compound in the liquid 
 (sodium bromide is much more soluble, and should be used, if 
 available). Warm, if necessary. Observe the odor (?), and 
 
 insert a rod dipped in ammonium 
 
 hydroxide solution (?) [Note 33, 
 
 p. 31]. 
 
 For the action of sulphuric 
 
 acid with a bromide, see 62 b. 
 b. Fit up a 250 c.c. flask with 
 
 a dropping-funnel and exit tube, 
 
 and connect with a U-tube (Fig. 
 
 13). Render the apparatus air- 
 Fig. 13 tight. Fill the U-tube with 
 
 dry, broken glass or porcelain, 
 mixed with a little red phosphorus (why?). Connect the 
 other limb of the U-tube with a second, larger U-tube [Store- 
 room] containing about 10 c.c. of water. Place about 5 g. 
 of red phosphorus mixed with twice its weight of sand in 
 the flask, add 5 c.c. of water, and mix by shaking. Pour 
 into the globe of the funnel about 8 c.c. of bromine [EXTREME 
 CARE. Do not spill upon the hands (Note 16, p. 2)]. Allow 
 the bromine to flow drop by drop on to the phosphorus, and 
 
47] BROMINE, IODINE, FLUORINE 43 
 
 let the gas dissolve in the water in the second U-tube. A 
 large volume of air is expelled before the hydrogen bromide 
 reaches the second U-tube. Disconnect the second U-tube 
 and reserve the solution for use in 46. 
 
 Try the effect of moist air upon the gas issuing from the 
 main apparatus (?). Hold in the gas a rod dipped in ammo- 
 nium hydroxide (?). 
 
 45. Properties of Aqueous Hydrobromic Acid. Divide the 
 solution into seven portions and examine its behavior toward 
 (a) litmus (?), (6) zinc in contact with a platinum wire (?), 
 (c) silver nitrate solution (?), (d) mercurous nitrate solu- 
 tion (?), (e) lead nitrate solution (?), (/) powdered manganese 
 dioxide (warm) (?). Boil c, d, and e, after pouring away the 
 supernatant liquid and adding more water to each (?). Com- 
 pare these results with those found in the case of hydrochloric 
 acid (32). 
 
 g. To the seventh portion add a few drops of chlorine-water 
 (43 6), a few drops of carbon disulphide, and shake (?). This 
 result indicates the relative affinities of chlorine and bromine 
 for what elements, and how? 
 
 46. Preparation of Iodine. Prepare a mixture of potassium 
 iodide (1 g.) and manganese dioxide (2 g.) exactly as in 42, 
 place it in an evaporating-dish, and moisten (2-3 drops) with 
 sulphuric acid diluted with water (1 Aq.: 2 acid). Cover the 
 dish with a watch-glass (partially filled with cold water to cool 
 the surface presented to the vapors) and warm very gently. 
 After a time examine the sublimate [Note 34, below]. Expose 
 a strip of filter paper, moistened with starch emulsion alone, 
 to the vapors (?). Recall the interaction of chlorides with 
 manganese dioxide and sulphuric acid [R 172]. 
 
 What other materials might be substituted for potassium 
 iodide in preparing iodine? 
 
 Note 34. Stains upon the fingers caused by iodine may be 
 removed by rubbing with sodium thiosulphate solution ("hypo"). 
 
 47. Properties of Iodine. Shake a few crystals of iodine 
 vigorously [Note 34] with about 10 c.c. of water in a test-tube. 
 Pour the clear liquid off, dividing it equally between four 
 test-tubes. 
 
 a. Add to one portion a few drops of chloroform, to another 
 ether, and to a third carbon disulphide. Shake each vigor- 
 ously and note the relative solubility (estimated by depth of 
 color) of iodine in water as compared with that in each of the 
 other solvents (?). 
 
44 BROMINE, IODINE, FLUORINE [ 48 
 
 b. Take 15 c.c. of starch emulsion and add the fourth portion 
 to it (?). Pour the mixture info the graduated cylinder and 
 add water so long as a sample poured out into a test-tube 
 continues to show an easily perceptible color. Why is the use 
 of starch considered to be a delicate test for iodine? Does it 
 show the presence of iodine in combination (48 6) ? 
 
 c. Repeat 43 b, using a very small crystal of potassium 
 iodide (?). Infer the relative activities of chlorine and iodine. 
 The result measures the relative affinities of these two halogens 
 for what elements? 
 
 d. Repeat c, using potassium iodide, but substituting bro- 
 mine-water for chlorine. In what order do the halogens stand, 
 in respect to activity, and why do you place them in that 
 order? 
 
 48. Preparation of Hydrogen Iodide [Hood]. 
 
 a. Repeat 44 a with potassium iodide or sodium iodide (?). 
 
 b. Use the apparatus in Fig. 13. Place in the flask a mix- 
 ture of finely powdered iodine (20 g.) and red phosphorus 
 (3 g.) intimately mixed in the mortar. Charge the U-tubes as 
 in 44 6. Place a little water in the dropping-funnel (or sub- 
 stitute, 36 6), warm the materials in the flask very slightly 
 (EXTREME CAUTION!] by waving the Bunsen flame once or 
 twice under the vessel, and allow the water to drop very slowly 
 upon them (?). After the air has all been expelled, and the 
 solution in the second U-tube has become sufficiently concen- 
 trated, remove the second U-tube. Test the issuing gas with 
 moist air (?) and with ammonia (?) as in 44 b. Hold in the 
 gas a piece of filter paper dipped in starch emulsion alone (?). 
 Explain. Reserve the contents of the second U-tube for use in 
 60 a. 
 
 49. Preparation of Hydriodic Acid [Hood]. Place 5 g. of 
 powdered iodine with 50 c.c. of water in a small flask provided 
 with a cork and an L-tube extending to the bottom. Pass 
 hydrogen sulphide from a Kipp's generator, or from the labora- 
 tory supply, through the mixture, loosening the cork once or 
 twice at first to permit the air to be displaced by the gas, until 
 the iodine is all gone and the solution no longer becomes brown 
 on being shaken. Agitate constantly to hasten the process. 
 Describe what happens. Warm and filter the solution. 
 
 Obtain a distilling-flask and condenser [Storeroom] and distil 
 the filtvate fractionally (Fig. 14), collecting first the part that 
 comes over at 100, then the parts boiling between 100-103, 
 103-106, and so forth. Use a very small flame, and be careful 
 not to allow it to reach the walls of the flask above the liquid, 
 or breakage will take place. A large flame may not only crack 
 
50] 
 
 BROMINE, IODINE, FLUORINE 
 
 45 
 
 the flask, but may also cause the thermometer to show a higher 
 temperature than it could acquire from the vapor alone. Stop 
 when the liquid is nearly all distilled off. Note the highest 
 
 Fig. I 4 
 
 temperature reached. Pour the residue into a test-tube and 
 keep the series for use in 50 b. 
 
 What substance causes the color of the higher fractions and 
 of the residue? Confirm your conclusion by a suitable test 
 (47). How is this colored substance formed? 
 
 50. Properties of Hydriodic Acid. 
 
 a. If 48 b was done, carry out the same experiments with 
 the solution as were made with the solution of hydrobromic 
 acid in 45 (?). Compare the results with those of 45 (?). 
 
 b. If 49 was done, add silver nitrate solution to each of the 
 fractions above 100 obtained in 49, using only a part of the 
 liquid in the case of the two with the highest boiling-points. 
 At what temperature did the most concentrated solution of 
 hydrogen iodide come over? What peculiarity of aqueous 
 hydriodic acid does the result indicate? What other solutions 
 show the same peculiarity [R 182, 232, 239, 241, 387, 440]? 
 
 Place a piece of zinc, in contact with a platinum wire, in the 
 remainder of one of the higher fractions (?). Test the other 
 with litmus paper (?), and then add pulverized manganese 
 dioxide, and warm (?). 
 
46 BROMINE, IODINE, FLUORINE [ 51 
 
 61. Hydrogen Fluoride. Cover a square of glass with a 
 thin layer of paraffin by warming it very cautiously far above 
 a Buusen flame and rubbing it on one side with solid paraffin. 
 Moisten about 3 g. of fluorspar in a leaden dish [Storeroom] 
 with concentrated sulphuric acid (do not cover with the acid). 
 With the end of a file draw some design upon the paraffin- 
 coated side, thus exposing parts of the glass to the action of 
 the vapor. Now cover the leaden dish with the glass, paraffin 
 side down, and set it in a moderately warm place, but not so 
 warm that paraffin may be likely to melt. After half an hour 
 or more remove the glass cover, warm, and wipe off the melted 
 paraffin with filter paper (?). Write equations representing 
 the action, and state what becomes of each of the constituents 
 of the glass [R 243, 521]. Try the test of a rod dipped in am- 
 monium hydroxide and held over the contents of the lead 
 dish (?). Does the gas fume with moist air? What sub- 
 stances, beside fluorspar, would serve the purpose of this ex- 
 periment? Why could not hydrochloric acid or nitric acid be 
 substituted here for sulphuric acid? What acid that we have 
 employed could be used here? 
 
 How may fluorine be liberated from a fluoride? Why can 
 it not be isolated from fluorides by the action of oxidizing 
 agents, as was the case with the other halogens? 
 
 52. Reducing Action of Hydrogen Iodide and Hydrogen 
 Bromide [Hood]. In connection with this experiment, it must 
 be kept in mind that an odor similar to that of rotten eggs 
 shows the presence of hydrogen sulphide (49), and an odor of 
 burning sulphur the presence of sulphur dioxide (13 a). 
 
 a. Pulverize finely about 1 g. of potassium iodide, place it 
 in a test-tube, and moisten with one or two drops of concen- 
 trated sulphuric acid (?). If too much acid has been taken, 
 start again. Warm gently. Investigate the result, which 
 furnishes a mixture of gases, as follows : 
 
 a. Breathe across the mouth of the test-tube to ascertain 
 the effect of the gas on moist air (?). What gases previously 
 made showed the same behavior? What do you infer in this 
 case? To confirm this conclusion, lower a glass rod dipped in 
 ammonium hydroxide into the test-tube (?). 
 
 ft. Is any characteristically colored vapor (?) mixed with 
 the gas recognized in a? Can you observe any other property 
 which identifies the substance? By what kind of chemical 
 action could this colored substance be formed from the product 
 identified in a? By what name is such a reaction known? 
 Was there any corresponding product formed when sulphuric 
 acid acted upon a chloride? 
 
53] BROMINE, IODINE, FLUORINE 47 
 
 7. Can you recognize still another (gaseous) product by its 
 odor? 
 
 The work in a and and 7 leads to the recognition of three 
 gaseous or vaporous products. Do not attempt to put all of 
 these in one equation. Construct an equation for the forma- 
 tion from the original materials of the gas recognized in a 
 (primary action), and make a separate equation for the forma- 
 tion of the other two products from the interaction of sulphuric 
 acid with the gas recognized in a (secondary action). What 
 two properties of sulphuric acid and what property of hydrogen 
 iodide are illustrated by this set of observations? 
 
 In case the above directions are not followed implicitly, and 
 large pieces of potassium iodide are taken, or too much sul- 
 phuric acid is used, still another gas (sulphur dioxide) may be 
 formed along with or instead of one of the above, and, in addi- 
 tion, a sublimate of free sulphur may be seen on the tube 
 [R 237]. 
 
 b. Repeat the work in a, using powdered potassium bromide 
 instead of the iodide, and answer the same questions. 
 
 63. Identification of Halogen Compounds. Imagine that 
 there are given to you four white substances, and that you know 
 them to be the fluoride, chloride, bromide, and iodide of some 
 metal. State what experiments you would make, and what 
 reasoning you would use, in order positively to identify the 
 halogen constituents of each. In two of these cases, two dif- 
 ferent actions have been encountered in this chapter and might 
 be used, and in the other two cases only one. Still another kind 
 of action might readily be thought of [R 63, 311]. Negative 
 results, say by showing that one is not a chloride, bromide, or 
 iodide, and is therefore a fluoride, must be confirmed by a 
 positive experimental test. 
 
 If the four hydrogen halides were given you in gaseous con- 
 dition in four jars, how should you proceed by chemical means 
 to identify each? 
 
CHAPTER IX. 
 
 DOUBLE DECOMPOSITION. OXYGEN COMPOUNDS OF THE 
 HALOGENS. HYDROGEN PEROXIDE. 
 
 64. Radicals and Double Decomposition. 
 
 a. To a few drops of potassium chloride solution add silver 
 nitrate solution (?). What kind of chemical interaction do 
 salts usually show in solution [R 264. See Note 4, p. 1] ? 
 Write the equation for this action. How can we tell whether 
 the precipitate is silver chloride, or potassium nitrate, or both? 
 Which is it (Appendix IV)? What was the interaction of 
 silver nitrate with hydrochloric acid (32 #)? To a few drops 
 each of solutions of two other chlorides, such as ferric chloride 
 and calcium chloride, add a little silver nitrate solution (?). 
 To find out whether all substances containing chlorine give 
 silver chloride in this way, try a few drops of potassium chlorate 
 solution with the same silver compound (?). State now what 
 radical a substance must contain in order that, with silver 
 nitrate, it may yield silver chloride (?). 
 
 b. To a few drops of silver sulphate solution add a solution 
 of any chloride (?). To find out whether all substances con- 
 taining silver yield silver chloride in this way, take a few drops 
 of silver nitrate solution in each of two test-tubes. To the 
 one portion add some ammonium hydroxide, and so obtain a 
 solution of ammonio-silver nitrate (Ag(NH 3 ) 2 N0 3 ). To the 
 other add some potassium cyanide solution [CAUTION! POISON!] 
 until the liquid is clear [Note 35, below], and so obtain a solu- 
 tion of potassium argenticyanide (KAg(CN) 2 ). Now add to 
 each of these a solution of sodium chloride (?) . Is the silver 
 radical present? Do all substances containing silver, when 
 mixed with a chloride, give silver chloride? Which compounds 
 alone give silver chloride by double decomposition? 
 
 c. Which substances alone will, by addition of mercurous 
 nitrate, give mercurous bromide (45 d)? Which substances 
 alone will, by addition of silver nitrate, give silver iodide (60 a 
 and 6)? Which substances alone will, by addition of an acid, 
 give hydrogen chloride (31)? Name the classes of substances 
 which are composed of radicals, and commonly interact by 
 double decomposition. 
 
 Note 35. An insoluble body will not dissolve as such merely 
 because of the addition of an excess of the precipitant, or even 
 
 48 
 
56] DOUBLE DECOMPOSITION 49 
 
 because of the introduction of a different reagent. When an 
 insoluble body appears to go into solution, the phenomenon indi- 
 cates that the substance added has interacted chemically with the 
 insoluble substance and has produced a new substance which is 
 soluble. 
 
 65. Chemical Equilibrium in Double Decomposition. 
 
 a. When solutions of two substances, each composed of two 
 radicals, are mixed, and no precipitate is observed, interaction 
 nevertheless occurs. Why was no precipitate observed when 
 solutions of silver nitrate and potassium chlorate (64 a) were 
 mixed? To answer this question, write the equation for the 
 double decomposition which might occur, and consider the 
 solubilities of the products (Appendix IV). A similar case 
 where the presence of the products is easily shown may now 
 be studied (see 6). 
 
 6. Place 10 c.c. of water in each of two test-tubes, add to 
 one a single drop of ammonium thiocyanate solution, and to 
 the other a single drop of ferric chloride solution. Now mix 
 the solutions (?). Write the equation for the action which 
 may be assumed to have occurred. Is there any evidence that 
 interaction has taken place? Which of the four is the colored 
 substance? Use the mixture for c. 
 
 c. When no precipitate is formed, is an action like the above 
 (a or 6) complete? To answer this question, divide the mix- 
 ture from 6 equally between four test-tubes. Keep one for 
 reference. To the second add one drop of ferric chloride solu- 
 tion (?), and to the third a drop of ammonium thiocyanate 
 solution (?). Interpret the result. Now add to the fourth 
 tube a few drops of ammonium chloride solution (?) and explain. 
 
 What other action have we shown to be reversible (33) ? All 
 double decompositions of substances composed of radicals are 
 reversible, like these two. They are also often far from com- 
 plete, when, as in the present instance, precipitation does not 
 occur. Why does precipitation tend to make the action more 
 nearly complete? 
 
 66. Hypochlorous Acid and Hypochlorites [Hood]. Fit up 
 a chlorine apparatus (29 6) capable of delivering a large amount 
 of chlorine and make sure that it is air-tight. Use the same 
 source of chlorine in 56 and 57. Between experiments, immerse 
 the end of the exit tube in sodium hydroxide solution, and 
 allow none of the gas to escape into the room. 
 
 a. Make an aqueous solution of chlorine in a ^test-tube. 
 Retain a few drops of this for use in 6, and place hi the re- 
 mainder some litmus paper, paper with printing [R 476] and 
 pen [R 754] and pencil [R 475] marks upon it, and a piece of 
 
50 DOUBLE DECOMPOSITION [57 
 
 colored calico. Observe the effect on each. Explain [R 176. 
 269]. 
 
 6. To a few drops of chlorine-water add a drop of indigo 
 solution [R 269] (?). 
 
 c. Place about 5 g. of quicklime in a small beaker, add a 
 few c.c. of water, warm slightly, and allow to slake (?). Now 
 add a little more water to make a thin paste, and pass chlorine 
 into the mixture for 10-15 minutes (?), keeping the vessel cool 
 by surrounding it with cold water (why?) and stirring the con- 
 tents during the process. Filter the paste, with the addition 
 of some water if necessary, and soak a piece of colored calico (?) 
 and some litmus paper in the filtrate. Remove these articles 
 to a beaker containing a little dilute sulphuric acid (?). Re- 
 peat these two operations with the same litmus and calico, if 
 at first little effect is seen. What substance produces the 
 effect? Why is the sulphuric acid required (see 55 a)? 
 
 What evidence does this experiment furnish that hypo- 
 chlorous acid is a more active oxidizing agent than is atmos- 
 pheric oxygen? Why is it thus more active? 
 
 How could you prepare an aqueous solution of pure calcium 
 hypochlorite [R 267, 268]? 
 
 67. Chlorates [Hood]. 
 
 a. Dissolve 3 g. (weighed on laboratory scales) of solid 
 potassium hydroxide in 7 c.c. of water in a test-tube and 
 saturate (Test? The solution must cease to feel soapy) the 
 solution with chlorine. While the saturation is proceeding, 
 calculate the volume of chlorine (at and 760 mm.) required 
 to interact with 3 g. of potassium hydroxide, one atomic 
 weight of chlorine being needed for each molecule of potassium 
 hydroxide. At five bubbles to 1 c.c., how many bubbles of 
 chlorine will be used? Observe how many bubbles issue from 
 your apparatus in fifteen seconds, and calculate how long the 
 operation may be expected to take (?). Crystals will appear 
 during the process of saturation and will increase in quantity as 
 the liquid afterwards cools. Filter off the crystals on a small 
 filter paper, and examine the filtrate and the crystals (in 6) 
 separately as follows : 
 
 Add to the filtrate dilute nitric acid (this is to destroy potas- 
 sium hydroxide, in case any remains : no equation needed) , and 
 test with a few drops of silver nitrate solution (?). What 
 radical is shown by this test to be present (54 a)? What 
 product is thus shown to have been formed by the interaction 
 of chlorine and potassium hydroxide? 
 
 6. Examine the crystals from a with a lens and describe 
 them. Dry the crystals, heat them in a narrow tube, and 
 
58] DOUBLE DECOMPOSITION 51 
 
 test for oxygen (?). Dissolve the residue from this operation 
 in distilled water and add silver nitrate solution (?). What 
 substances constituted the crystals and the residue, respec- 
 tively? From the behavior of the former substance during 
 making, what do you infer as to its solubility? Is the infer- 
 ence correct (Appendix IV) ? 
 
 What effect was observed on adding silver nitrate solution to 
 a solution of potassium chlorate (54 a)? How may the chlo- 
 rate radical be distinguished from that of the chlorides? The 
 crystals of potassium chlorate made in a are not free from 
 traces of potassium chloride (why?), and could not therefore 
 be utilized for this test. What method' should you suggest 
 for purifying the chlorate [R 273]? 
 
 To a minute amount of finely powdered potassium chlorate 
 add a few drops of pure, concentrated hydrochloric acid 
 (66 a) (?). The yellow substance is formed by decomposition 
 of one of the products [R 275] (?). How would a chloride 
 behave with hydrochloric acid? 
 
 Give the three ways of distinguishing chlorides from chlo- 
 rates. 
 
 68. Perchlorates. 
 
 a. Measure 600 c.c. of water into your 1-liter bottle, and 
 mark the level reached. Observe the temperature and pressure 
 of the air and calculate the weight of potassium chlorate which 
 will be necessary to give 600 c.c. of oxygen under these conditions 
 (the tension of aqueous vapor may be neglected, as the vapor 
 will occupy only about 10 c.c. of the 600 c.c. at 18) and at the 
 same time leave the perchlorate and chloride as a residue. 
 This stage is reached when one-fifth of the total oxygen has been 
 evolved : in other words, take so much of the chlorate as, if 
 completely decomposed, would furnish five times 600 c.c. 
 
 Fill the 1-liter bottle with water and invert it over the pneu- 
 matic trough. Weigh the calculated amount of chlorate into 
 a hard glass test-tube, which has previously been closely fitted 
 with a one-hole cork and delivery tube, and see that the appa- 
 ratus has been made air-tight. Gently heat the chlorate and 
 collect in the 1-liter bottle enough oxygen to fill the bottle to 
 the mark, measured, of course, when the mark is at the same 
 level as the water in the trough. Proceed slowly towards the 
 end so as to allow the gas to cool, stop heating when the mark 
 is reached, and remove the delivery tube at once from the 
 water. Pour the melted substance into a mortar before it has 
 time to solidify. Pulverize the mixture. 
 
 The mixture consists mainly of the chloride and perchlorate 
 
52 DOUBLE DECOMPOSITION [59 
 
 of potassium. The solubilities (grams of the salt dissolved by 
 100 c.c. of water) of these salts are as follows: 
 
 15 20 100 
 
 Potassium chloride 33 35 56 
 
 Potassium perchlorate ... 1.5 1.8 20 
 
 To separate the substances, calculate approximately the amount 
 of potassium chloride which must be present, and shake the 
 powder persistently with an amount of cold water just suffi- 
 cient to dissolve this salt. Cut, fold, and place in a funnel a 
 filter paper just large enough to hold the undissolved material. 
 Collect the latter upon the filter and wash it with a few drops 
 of cold water. Calculate the amount of water which at 100 
 will dissolve the residue, assuming it to be potassium perchlo- 
 rate. Dissolve it in this amount of water by boiling, and allow 
 the solution to stand for an hour or two. Collect the crystals 
 upon a filter, wash them as before, and dry them on a radiator. 
 
 6. Dissolve a little of the substance in distilled water and 
 test with silver nitrate solution (?). Explain (66 a). 
 
 To a minute amount of the crystals add a few drops of pure 
 concentrated hydrochloric acid [R 276] (?). Why does the 
 result differ from that when potassium chlorate was treated 
 with the same acid (67 6)? 
 
 Place about 1 g. of the crystals in a narrow test-tube, heat, 
 and test for oxygen (?). 
 
 How could you distinguish a perchlorate from a chloride, and 
 from a chlorate? 
 
 69. Bromic and lodic Acids. 
 
 a. Take two test-tubes and place in one a minute fragment 
 of iodine and in the other 2-3 drops of bromine-water. Add 
 about 5-10 c.c. of water and a little carbon disulphide to each, 
 and shake (43 a and 47 a). The carbon disulphide is added 
 simply for the purpose of collecting the halogen and making its 
 
 Eresence obvious. Now pass chlorine (generated as in 43 6) , a 
 iw bubbles at a time, through (or add chlorine-water, a few 
 drops at a time, to) the water in the test-tubes, alternately, 
 and shake vigorously after each addition of chlorine until a 
 change is seen [R 277] (?). Which of the halogens is first 
 affected, and why? 
 
 b. To about 10 c.c. of water in a test-tube add a single drop 
 of potassium iodide solution and then a single drop of potas- 
 sium bromide solution. Introduce also a few drops of carbon 
 disulphide. Now pass chlorine (generated as in 43 6), a few 
 bubbles at a time, into the liquid, or add chlorine-water a 
 
60] - DOUBLE DECOMPOSITION 58 
 
 little at a time, shaking vigorously after each addition of 
 chlorine (?). Continue until no further changes occur, and 
 explain all the changes which are observed. This procedure 
 is used in analysis for recognizing a bromide in presence of an 
 iodide. 
 
 c. Prepare some dilute bromic acid by taking 2-3 c.c. of 
 potassium bromate solution and adding an equal volume of 
 dilute sulphuric acid (?). Make the equation for this double 
 decomposition. Is the action complete (55)? In what fol- 
 lows, disregard the substances present with the bromic acid, 
 and place the latter only in the equation. Drop into this solu- 
 tion a single small crystal of iodine and shake repeatedly, 
 allowing the mixture to stand for some minutes after each shak- 
 ing^). Pour off the solution from any undissolved iodine, and 
 to the clear liquid add a few drops of carbon disulphide (?). 
 What free halogen is here detected? What does this show in 
 regard to the relative tendencies of bromine and iodine to unite 
 with oxygen? What inference can you draw in regard to the 
 relative activity of chlorine towards oxygen? What would 
 be the action of iodine upon a solution of chloric acid? 
 
 Which variety of chemical change was here observed? How 
 could you show that, although no change is visible, the bromic 
 acid actually is formed when the sulphuric acid is added, above, 
 and that it is the bromic acid, and not the potassium bromate, 
 which interacts with the iodine? 
 
 60. Peroxides. 
 
 a. Dissolve about 2 g. of sodium peroxide in 100 c.c. of cold 
 water in a flask. Add this amount of the oxide, a very little at 
 a time, shaking and cooling (why? [R 571]) the mixture in a 
 stream of water during the process. While still cooling the 
 solution, add to it dilute sulphuric acid a few drops at a time 
 until the mixture is acid (test?). What does the liquid now 
 contain (55 a)? Divide the mixture into four parts and use 
 them in b, c, d, and e. 
 
 b. To one portion, contained in a small test-tube, add finely 
 powdered manganese dioxide (?). Test the escaping gas for 
 oxygen. What role does the manganese dioxide play here? 
 
 c. Prepare a solution containing free permanganic acid and 
 sulphuric acid by adding a large excess of dilute sulphuric 
 acid to 5 c.c. of potassium permanganate solution (equation? 
 See 55). Add some of this mixture to the second portion 
 (from a). Test for oxygen the gas which comes off (?). What 
 variety of chemical activity does the hydrogen peroxide show 
 here? 
 
 d. To the third portion add some starch emulsion containing 
 
54 DOUBLE DECOMPOSITION [ 60 
 
 a drop of potassium iodide solution (?). In writing the equa- 
 tion for this action remember that the solution of hydrogen 
 peroxide from a contained excess of sulphuric acid, which will 
 interact (65) with the potassium iodide (?). The product of 
 this action then interacts with the hydrogen peroxide. What 
 variety of chemical activity does the hydrogen peroxide show 
 here? 
 
 e. To the fourth portion add 5 c.c. of ether (object of this 
 [R 306]?) and shake, and then add one drop and no more of 
 potassium dichromate solution and shake again (?). This is 
 one of the most characteristic and delicate tests for hydrogen 
 peroxide. 
 
 State what substance here combines with the hydrogen per- 
 oxide and how it is formed [R 305-306]. Make no equation. 
 
 /. Suspend lead dioxide, barium dioxide, and pulverized 
 manganese dioxide, separately, in water, add dilute sulphuric 
 acid and shake for some time, cooling as in a. Filter, and apply 
 to each filtrate the test described in e (?). What are the dif- 
 ferences in behavior and constitution between a true peroxide 
 and those oxides which are sometimes called peroxides [R 308] ? 
 
 g. What are the radicals of: Bromic acid, potassium chlorate, 
 hypochlorous acid, bleaching powder, sodium peroxide, potas- 
 sium permanganate, iodic acid, potassium perchlorate? 
 
CHAPTER X. 
 
 IONIZATION AND INTERACTIONS OF ACIDS, BASES, AND 
 SALTS. 
 
 61. lonization. Name the four distinct methods by which 
 we may ascertain experimentally whether a substance is 
 ionized in solution or not, and learn the extent of the ioniza- 
 tion [R 289, 292, 293, 328]. Define the term ionization. 
 
 The degrees to which aqueous solutions of many substances 
 are ionized are given in Appendix VI. Constant reference to 
 this will be necessary in interpreting the observations in this and 
 succeeding chapters. 
 
 The experiments of this paragraph may be postponed until 
 after the work in 62 or 63 has been done, if a set of the appa- 
 ratus is not available at this moment. 
 
 Obtain [Storeroom] a pair of electrolytic cells* (Fig. 15) and 
 
 Fig. is 
 
 half fill one with dilute sulphuric acid. When some material 
 has been placed in the second cell, and both cells are connected 
 
 * Each cell consists of a glass, flat-bottomed, specimen tube (about 
 75 x 22 mm.) fitted with a two-hole rubber stopper in which a ver- 
 tical groove has been cut to permit the escape of gases. The electrodes 
 are pieces of glass tubing (about 10 cm. long) into which platinum 
 wires have been sealed by means of sealing-glass. Contact is made by 
 means of a thin copper wire welded to the inner end of the platinum 
 wire. The platinum wire projects about 15 mm. and is turned upward 
 and stuck to the outside wall of its tube by means of sealing-glass. To 
 
 66 
 
56 
 
 ACIDS, BASES, AND SALTS 
 
 [ 61 
 
 in series with the source of electricity,* an evolution of gas in 
 the first cell will indicate that the circuit is complete and that, 
 therefore, the material which has been placed in the second cell 
 is a conductor. If, on the other hand, the material in the second 
 cell is a non-conductor, or even a very bad conductor, no evo- 
 lution of gas will be observed in the first cell. If the material 
 in the second cell is a solution and a conductor, what conclusion 
 may be drawn in regard to the condition of the dissolved body 
 [R 326] ? 
 
 Half fill the second cell with the substances named below in 
 turn. See very particularly that the electrodes in each cell are 
 not touching one another. Connect with the battery, and ob- 
 serve the effect in the first cell. When the same experiment 
 has been shown in the class-room, the result may be recorded 
 here and the experiment omitted. Wash the second cell and 
 electrodes very carefully after each trial. 
 
 The following eight substances, or solutions, show the behavior 
 typical of the classes of materials to which each example belongs. 
 After giving the result in your notes, name the class which is 
 illustrated in each case. 
 
 Fig. 16 
 
 secure a large electrolytic surface, the tubes are dipped for a distance 
 of 20 mm. in strong chloroplatinic acid solution and heated in the 
 Bunsen flame. This leaves a coating of metallic platinum on their 
 surface. 
 
 The cells are set into holes bored in a small block of wood. To 
 protect the latter from the action of acids the blocks should be prepared 
 by soaking in hot paraffin. 
 
 * A storage battery of three cells in series may be used, but is 
 always in danger of being ruined by short-circuiting through careless- 
 ness. Protection by means of a fuse leads to continual interruptions 
 of the work. The best plan is to use the current passing from the 
 lighting circuit (D, Fig. 16), through two 100-candle-power lamps (or 
 an equivalent arrangement of other lamps), one on each wire from the 
 dynamo, and to reduce the voltage to 6-8 volts by means of a suitable 
 shunt (S). With a lamp on only one wire, there is danger that a student 
 may produce a short circuit by allowing the other wire to touch a gas 
 connection or steam pipe. 
 
62] ACIDS, BASES, AND SALTS 67 
 
 a. Dry, crystallized sodium chloride (?). 
 6. Distilled water (?). 
 
 c. Aqueous solution of sodium chloride [Side-shelf] (?). 
 
 d. Diluted aqueous solution of sodium hydroxide [Desk] (?). 
 
 e. Diluted aqueous solution of hydrogen chloride [Desk] (?). 
 /. Aqueous solution of sugar (?). Dry the cell by washing 
 
 first with alcohol and then with ether. 
 
 g. Toluene in the dried cell (?). 
 
 h. Hydrogen chloride dissolved in dry toluene [Side-shelf] (?). 
 What difference between water and toluene do e and h bring to 
 light? Keep this solution corked up in a dry test-tube for use 
 in 65 d. 
 
 62. Ionic Materials. The chemical composition of the ions 
 into which any compound is divided by solution in water may 
 be ascertained in two ways: (1) by electrolysis and examina- 
 tion of the substances liberated at the electrodes [R 310-312], 
 and (2) by studying the interactions (particularly the double 
 decompositions, 64) of the substance with other substances 
 [R 281-283], for the radicals and ions of a substance are the 
 same. What classes of chemical compounds are alone ionized? 
 
 The ionic substances may be named as follows: 
 
 Substance. Name. 
 
 Ionic sodium (Na*) Sodium-ion. 
 
 Ionic hydrogen (H*) Hydrogen-ion. 
 
 Ionic chlorine (Cl') Chloride-ion. 
 
 Ionic chlorate radical (CKV) Chlorate-ion. 
 
 Ionic hydroxyl (OH') Hydroxide-ion. 
 
 Ionic ferric iron (Fe"') Ferric-ion. 
 
 Ionic ferrous iron (Fe") Ferrous-ion. 
 
 Ionic bisulphate radical (HSO/) Hydrogen-sulphate-ion. 
 
 In using these terms, note that sodium-ion (with the hyphen) 
 is the name of the substance, and not of the hypothetical, 
 charged atom. When speaking in terms of hypothesis, there- 
 fore, we may not say "a sodium-ion," or "sodium-ions," any 
 more than we should say "an ionic sodium" or "ionic sodiums." 
 To describe the charged atom or group of atoms, we must write 
 "a sodium ion," "sodium ions," "chlorate ions," etc. 
 
 In accordance with the above nomenclature, give the names 
 and symbols (not forgetting the charges) of the three chief 
 physical components, i.e., distinct substances (in addition to 
 water), present in the aqueous solutions of sodium chloride, 
 hydrogen chloride, and sodium hydroxide [R 334] (?). In 
 
68 ACIDS, BASES, AND SALTS [ 63 
 
 these solutions there are still other physical components pres-. 
 ent in small amounts which are these in each of the cases 
 just mentioned [R 344]? Give a concise comparative state- 
 ment of the specific physical properties (such as, color, molec- 
 ular weight, solubility, behavior towards electrically charged 
 bodies, physical state) of the ionic and the free forms of sodium, 
 hydrogen, and chlorine (?). 
 
 63. Relations of the Molecular Substance to its Constit- 
 uent Ionic Substances (in Equilibrium). The ions of an ion- 
 ogen and the remaining molecules are in chemical equilibrium 
 [R 297]. What changes take place, respectively, when the 
 solution is concentrated by evaporation and when it is diluted, 
 as in a, below [R 297-298, 335] ? Can the proportion of mole- 
 cules be increased otherwise than by concentrating the solu- 
 tion (see 6, below) [R 335-336] ? With substances like sodium 
 chloride and hydrogen chloride these changes cannot be per- 
 ceived by the eye (why?). In the following instances (a and 
 b) the ionic and molecular substances are both perceptible to 
 the eye, and their relations as described above may, therefore, 
 be studied very easily. 
 
 a. Examine a solution of potassium bromide. What is the 
 color of bromide-ion? Take a minute amount (say 0.2 g.) of 
 cupric bromide in a dry test-tube. Add two drops of water and 
 agitate for some time (?). Then add more water, a drop or 
 two at a time, agitating vigorously, and giving the substance 
 time to dissolve, if it can, after each addition. Continue the 
 addition of water cautiously until the substance has all dis- 
 solved, and afterward until the change in color is complete, 
 and then stop. What is the color of the molecules of cupric 
 bromide? What is the color of cupric-ion? Compare the 
 color with that of cupric sulphate solution (?) and explain. 
 Formulate the change which has been witnessed. 
 
 b. Now take a fresh portion of cupric bromide and repeat 
 the experiment as in a, stopping the addition of water at the 
 green stage. Divide the mixture into two parts. To one add 
 2-3 g. of solid potassium bromide and shake vigorously (?). 
 To the other portion add 4-5 g. of solid cupric chloride (?). 
 Interpret the results. 
 
 The converse case, in which one of the ions is removed and 
 the dissociation is promoted, is discussed in d. 
 
 c. Many ionic substances are colored, although the color 
 does not always differ markedly from that of the molecules. 
 Examine the following solutions [Side-shelf], make a list of the 
 ionic substances contained in them, with their formulae and 
 charges, and note the color of each kind of ions : cobalt chloride, 
 
64] ACIDS, BASES, AND SALTS 59 
 
 potassium permanganate, potassium dichromate, chrome-alum 
 (this last solution freshly made by dissolving the solid). 
 
 d. Just as the union of ions to form molecules may be pro- 
 moted by addition of a substance yielding a common ion (? 
 63 b), so the reverse of this, namely, the dissociation of mole- 
 cules into ions, may be promoted by the removal of one of the 
 ionic materials. The removal of ions may be accomplished in 
 several ways [R 360-364]: 
 
 (1) By union of the ion with some other ion, when the mole- 
 cules thus formed are insoluble. This case will be illustrated 
 next (see 64). 
 
 (2) By union of the ion with some other ion, when the mole- 
 cules thus formed, although soluble, are very little dissociated 
 by water (see 66, 75, and 92 a). 
 
 (3) By discharge of the ion and liberation of its material, 
 through transfer of the charge to another substance (see 68). 
 
 (4) By union of the ion with some other material to form a 
 compound ion, illustrated in 54 b. 
 
 (5) By decomposition of the ion, as in 59 c. 
 
 (6) By the mere change in the valence (amount of the charge) 
 of an ion, for this converts it into another substance of the 
 same material composition (see 160 g, 167 b, 168 d). 
 
 (7) By discharge of the ion and liberation of its material 
 through electrolysis. This was illustrated in 61. 
 
 64. Precipitation on Mixing lonogens. 
 
 a. Place 3-4 c.c. of silver nitrate solution in a test-tube 
 and dilute with water. Add potassium chloride solution 
 cautiously and agitate continuously, until no further precipi- 
 tation occurs (?). Filter, concentrate the filtrate by evapora- 
 tion, and pour it into a watch-glass to crystallize (?). Two 
 salts are formed. 
 
 Formulate the action (Fig. 17). In doing this, show the 
 
 K- +C1' 
 
 Tl U 
 
 (dsolvd.) KNO 3 AgCl (dsolvd.) 
 
 U 
 
 AgCl (solid) 
 Fig. 17 
 
 three main physical components of each of the original solu- 
 tions and the relations of these components (in equilibrium) to 
 one another in each case. Show also the molecular products 
 
60 ACIDS, BASES, AND SALTS [ 65 
 
 formed when the solutions are mixed. Assuming that the 
 solutions are approximately normal [R 148], ascertain the 
 proportions in which the original components are present 
 before mixing (Appendix VI). Learn also to what extent the 
 molecular products will be formed by union of the ions (Appen- 
 dix VI), and, in the case of an insoluble substance, how com- 
 plete will be the precipitation (Appendix IV). On the basis 
 of this complete information, explain in detail, and one by one, 
 in what way, and to what extent, each of the six original com- 
 ponents is affected by the results of mixing. 
 
 Name the components of the filtrate and explain how each 
 is affected by the evaporation and crystallization. 
 
 How does the formation of the precipitate of silver chloride 
 illustrate 63 d (1)? Upon what factor does the completeness 
 of the change depend? Is, or is not, silver chloride a highly 
 ionized substance [R 332] ? 
 
 Aside from double decompositions, what means have we for 
 learning of what radicals a salt (like silver nitrate) is composed? 
 
 6. To a little cupric sulphate solution in a test-tube add 
 sodium hydroxide solution (?). Exactly as in 64 a (second 
 par.), formulate, study and explain the whole action. How 
 does this illustrate (1) of 63 d ? 
 
 To what classes of ionogens do the four molecular substances 
 respectively belong? 
 
 c. To a little cupric sulphate solution add a little dilute 
 hydrochloric acid (?). In what respects does the result differ 
 from those in a and b, and why? Can any acids be prepared 
 by precipitation, and if so, which? Give illustrations. 
 
 65. Bases and Acids: Hydroxide-ion and Hydrogen-ion: 
 Indicators. 
 
 a. Examine distilled water in respect to (a) taste, (b) be- 
 havior with litmus, (c) conductivity (done already, 61 6). 
 
 6. Dissolve a small piece of sodium hydroxide in water and 
 examine the solution in respect to (a) taste, by diluting a 
 little and tasting one drop, (b) behavior with litmus, (c) be- 
 havior with phenolphthalem, (d) conductivity (see 61 d). 
 These properties belong to aqueous solutions of all bases. 
 Aside from the water, what component alone is common to all 
 such solutions and has the above properties? 
 
 c. Examine an aqueous solution of hydrochloric acid in 
 respect to (a) taste, (b) behavior toward litmus, (c) behavior 
 with phenolphthalem, (d) conductivity (see 61 e), (e) action 
 on a piece of marble, (f) action on an iron nail (clean this with 
 the file before use). These properties are shown by all aqueous 
 solutions of acids. Aside from the water, what component 
 
66] ACIDS, BASES, AND SALTS 61 
 
 alone is common to all such solutions and has these proper- 
 ties ? 
 
 d. Take the solution of hydrogen chloride in toluene (61 H) 
 and examine it in respect to (a) conductivity (see 61 h), (b) 
 action on a piece of marble, dried in advance by heating in a 
 dry porcelain dish for a few moments, (c) action on an iron 
 nail (clean as before). Be sure that perfectly dry vessels are 
 used in these experiments. What substance identified in c is 
 absent from this solution? What difference between water 
 and toluene, as solvents, does this result indicate? In what 
 three other respects would the two solutions of hydrogen chlo- 
 ride be found to differ? 
 
 66. Ionic Chemical Changes: I. Union and Disunion of 
 Ions: Neutralization (Union of H* and OH'). (Two students 
 working together.) A considerable chemical change may occur 
 not only in precipitation (see 64), but also when ions unite to 
 form a substance which, although soluble, is very little ionized 
 by water. Such changes are illustrated in 66 and 67. 
 
 Weigh out 5 g. (laboratory scales) of potassium hydroxide, 
 dissolve it in 50 c.c. of distilled water, and pour the clear solu- 
 tion into a burette (Fig. 2, p. 7). Measure 10 c.c. of con- 
 centrated hydrochloric acid in the graduated cylinder, mix it 
 in a small beaker with 50 c.c. of distilled water, and pour this 
 into a second burette. Now, into a small beaker or flask run 
 15 c.c. of the acid solution from the second burette, and then 
 add to it two drops of phenolphthalein solution. Place the 
 vessel under the first burette, read the level of the liquid in 
 the burette, and allow the alkali to run into the acid drop by 
 drop, stirring constantly, until the last drop confers the faintest 
 perceptible pink tinge on the whole solution. If you do not at 
 first succeed in stopping at the right point, repeat the experi- 
 ment. Note the volume of alkali used. Concentrate the solu- 
 tion on the sand bath until a drop deposits crystals on cooling, 
 and then remove the dish from the sand bath promptly and 
 set it aside. 
 
 When sufficient crystals have appeared, dry them with filter 
 paper and examine with respect to (a) form, (b) taste, (c) 
 exposure to moist air, (d) action of a solution on litmus, (e) 
 conductivity of aqueous solution (done already, 61 c). Con- 
 struct a table comparing the substance in these respects with 
 the materials from which it was made. Compare the substance 
 with potassium chloride on the side-shelf. How should you 
 determine whether a substance obtained in this way contained 
 "water of crystallization" or not? Make the necessary experi- 
 ments (?). Wash out the burettes. 
 
62 ACIDS, BASES, AND SALTS [ 67 
 
 Following the directions in 64 a (second par.), formulate, 
 study, and explain the whole action. Note, however, that 
 there is here no insoluble substance. Show how this experi- 
 ment illustrates 63 d (2). 
 
 Express the change involved in every neutralization of highly 
 ionized substances by means of the simplest equation. How 
 may neutralization be defined, in terms of the hypothesis of 
 ions? How may it be denned, taking account of all the facts, 
 but omitting all reference to ions 
 
 Calculate the approximate concentration, in terms of a nor- 
 mal solution as unity, of the potassium hydroxide solution 
 used above (?). From the volumes of alkali and acid used in 
 neutralizing, calculate the concentrations of the diluted hydro- 
 chloric acid, and of the concentrated acid employed to make 
 it, expressing the concentrations in terms of a normal solution. 
 Calculate the number of grams of hydrogen chloride per liter in 
 the dilute and concentrated acids, respectively. 
 
 67. Neutralization of Slightly Ionized and of Insoluble 
 Substances. 
 
 a. Consider the degree of ionization of acetic acid (Appendix 
 VI). To neutralize 1 liter of normal acetic acid, would more 
 or less alkali be required than to neutralize 1 liter of normal 
 hydrochloric acid? In what way, precisely, would the details 
 of the change be different in the case of acetic acid [R 356]? 
 Name some of the consequences of this difference [R 357-358]. 
 
 b. Dilute a few drops of cupric sulphate solution with water 
 and add excess of sodium hydroxide solution (?). Fit a filter 
 paper properly into a funnel (Fig. 3, p. 9). Filter the mix- 
 ture, and wash the precipitate (?) repeatedly with distilled 
 water to remove soluble substances. Now place a clean test- 
 tube below the funnel, perforate the bottom of the filter paper, 
 and wash the precipitate through into the test-tube by means 
 of a stream of water from the wash-bottle. To the suspended 
 cupric hydroxide, cautiously add dilute hydrochloric acid in 
 amount just sufficient to give a clear liquid. Concentrate the 
 liquid on the sand bath until a drop, removed to a watch-glass, 
 shows signs of crystallizing when cold. Then remove the dish 
 promptly from the sand bath and allow it to cool. Examine 
 the crystals (?). 
 
 Formulate this action as in 64 a, taking account, however, 
 of the fact that one of the interacting substances is an "insol- 
 uble" solid [R 349]. Describe in detail the stages through 
 which the final production of solid cupric chloride is accom- 
 plished. 
 
 To what class of ionic chemical changes [R 360] does the 
 
68] ACIDS, BASES, AND SALTS 63 
 
 foregoing action belong? Answer the same question in regard 
 to the precipitations of salts in 64. 
 
 68. Ionic Chemical Changes: II. Displacement. 
 
 a. Place several pieces of granulated zinc in a dilute solution 
 of cupric sulphate and set aside until the change is complete 
 (test?). Occasional agitation will hasten the change (why?). 
 Filter. What is the precipitate [R 361]? Before examining 
 the filtrate, take a few drops of cupric sulphate solution and a 
 like amount of zinc sulphate solution in two test-tubes. Dilute 
 each solution with water, and add to each ammonium sulphide 
 solution (?). 'What is the precipitate in each case, and what 
 ions are required to form it? 
 
 To the filtrate from the first part of this experiment add 
 ammonium sulphide solution (?). What ions were present in 
 the filtrate? What changes did the zinc and the cupric ions, 
 respectively, undergo in the first part of the experiment? For- 
 mulate this change in an equation. In the course of this ex- 
 periment, what becomes of the molecular cupric sulphate 
 (63d (3))? 
 
 What substances could have been substituted for the cupric 
 sulphate without affecting the result? What substances, be- 
 side zinc, would have precipitated copper (Appendix VII) ? 
 What other elements, beside copper, are displaced by zinc? 
 
 Which of the elements displaced by zinc did we prepare in 
 quantity by an action like the present (18). Formulate this 
 action in terms of the hypothesis of ions. What were the 
 products of the action of zinc upon concentrated sulphuric 
 acid (16 d)1 What is the chief component of this form of 
 the acid (Appendix VI) ? If this component interacted with 
 the zinc, to what class of chemical changes did this action 
 belong? 
 
 6. Formulate the actions in 16 a in terms of the hypothesis 
 of ions. Explain in terms of the hypothesis the differences in 
 activity of the various metals [R 362]. 
 
 c. Examine your notes on Chap. VIII. Formulate the fol- 
 lowing actions in terms of the hypothesis of ions : 
 
 Free chlorine and bromide-ion (43 b and 45 g). 
 Free chlorine and iodide-ion (47 c and 50 a). 
 Free bromine and iodide-ion (47 d) 
 Free iodine and sulphide-ion (49). 
 
 Arrange these four elements in a series similar to the electro- 
 motive series of the metals (Appendix VII). Where should 
 you place fluorine in this series [R 241] ? Can you indicate the 
 approximate position of oxygen [R 239, 241]? 
 
64 ACIDS, BASES, AND SALTS [ 69 
 
 69. Ionic Chemical Changes: III, IV, V. In addition to 
 union or disunion of ions (I), and displacement (II), there are 
 three other ways in which ions may undergo chemical change. 
 
 One of these (III) was illustrated in 29 a and b, in 54 b, in 
 67 a and b, in 69 a, b, and c, and in 60 c (?). 
 
 Another (IV) was illustrated in 29 a (?). 
 
 Still another (V) was illustrated in 61 (?). 
 
 Define each of these three classes of ionic chemical change 
 and formulate the illustrations anew so as to show how the 
 actions cited belong to the class illustrated. 
 
 Re-examine the seven ways in which ionic substances are 
 removed, and dissociation of the parent molecules is promoted 
 (63 d), and indicate that one of the five classes of ionic change 
 to which each of the seven belongs. 
 
 70. Non-Ionic Actions. In previous class-room and lab- 
 oratory experiments we have observed the formation of iono- 
 gens in other ways than those illustrated in this chapter. These 
 ways are non-ionic, or not distinctly ionic. Give illustrations 
 of such of these ways as you recall : acids, two ways ; bases, one 
 way; salts, four ways, together with the reference numbers of 
 the laboratory experiments in which they occur. 
 
CHAPTER XL 
 
 SULPHUR. 
 
 71. Sulphur. 
 
 a. In a dry test-tube place a very small piece of roll sulphur 
 with 2-3 c.c. of carbon disulphide, and shake (?). Allow the 
 clear solution to evaporate spontaneously (i.e., without the aid 
 of heat) in a watch-glass [Hood], and describe the crystals 
 [R 138] (?). See Note 25, p. 10. 
 
 b. In a dry test-tube place about 5 g. of roll sulphur, melt 
 the substance with the least possible application of heat (the 
 material must remain pale-yellow), and pour the melt into a 
 beaker of cold water. Dry some of the product with filter 
 paper and test its solubility in carbon disulphide as in a. 
 
 c. Melt about 10 g. of sulphur in the same test-tube, and heat 
 the body until it boils (?). Note the changes in color and 
 fluidity that occur. To learn the nature of the substance 
 formed by heating, chill the sulphur while it is boiling vigor- 
 ously by pouring it suddenly into cold water (?). Note the 
 physical state of the product, dry a part with filter paper, and 
 examine its solubility in carbon disulphide (?). Set the re- 
 mainder aside for a few days, and then study its appearance 
 and solubility again (?). Keep also the test-tube from which 
 it was poured, and examine, at the same time, in both' these 
 respects, the sulphur which remained adhering to its walls and 
 was not cooled so suddenly (?). Account for the change when 
 sulphur is heated, and the differing results of rapid and slow 
 cooling [R 369]. 
 
 Why are we convinced that none of the changes was due to 
 interaction with the water? 
 
 d. Mix in a mortar 2 g. of iron filings and 1 g. of powdered 
 sulphur. Transfer to a dry test-tube and heat gently (?). 
 When cool, break the test-tube in a mortar and use the black 
 product (?) in 72 a. 
 
 Recall and record here a case of the union of sulphur with a 
 metal which was obseryed previously. 
 
 72. Hydrogen Sulphide [Hood]. 
 
 a. Place a part of the product from 71 d, or about 1 g. of 
 ferrous sulphide, in a test-tube and add a little dilute hydro- 
 chloric acid (?). Note the odor of the gas, and apply to it a 
 strip of filter paper dipped in lead nitrate solution [R 705] (?). 
 
 6. With a Kipp's apparatus generating hydrogen sulphide, 
 
 65 
 
66 SULPHUR [ 73 
 
 or with the laboratory supply of the gas, connect a glass nozzle. 
 When the air has been displaced, set fire to the gas. Describe 
 the color of the flame. Hold a porcelain dish in the middle of 
 the flame for a few moments (?). What substance is depos- 
 ited, and must, therefore, exist uncombined in the interior of 
 the flame? What other substance does this justify us in 
 assuming to be liberated in the same region? What do these 
 facts indicate regarding the stability of the gas when heated, 
 and the difficulty, therefore, of making the compound by the 
 direct union of its elements? What are the products of the 
 complete combustion of the gas, and in what two stages does 
 this combustion take place? Make equations showing both 
 stages. 
 
 73. Properties of Aqueous Hydrogen Sulphide : I [Hood]. 
 
 a. Take about 15 c.c. of water in a carefully cleaned test- 
 tube and saturate (test? Note 36, below) it with hydrogen 
 sulphide. Test the solution with litmus paper [R 347] (?). 
 Pour one-third or less of the solution into another test-tube 
 and boil this portion vigorously, noting from time to time the 
 odor of the vapor. Can this gas be driven off completely by 
 boiling? Does hydrogen chloride solution behave in the same 
 way or differently [R 182] ? 
 
 6. Divide the remainder of the solution into two parts and 
 add 2-3 g. of iron dust (pulverized iron) to one of them. Shake 
 vigorously and then allow the mixture to stand (?). After a 
 few minutes collect the insoluble matter upon a filter and wash 
 until it no longer smells of hydrogen sulphide. Transfer the 
 precipitate to a test-tube by puncturing the paper and washing 
 through. Add dilute hydrochloric acid to the product and 
 note the odor (?). What substance must have been present 
 in the precipitate, and how was it formed? Account for the 
 extreme slowness of the action of iron upon the solution of 
 hydrogen sulphide [R 374]. 
 
 c. Allow the third portion of the solution made in a to stand 
 for some days exposed to the air (?). Explain the turbidity. 
 
 d. If 49 was performed, record here the action which took 
 place. If 49 was not performed, place a single crystal of 
 iodine in 5 c.c. of water and saturate (test? Note 36, below) the 
 liquid with hydrogen sulphide [R 238] (?). 
 
 What ionic substance is shown by the first part of 73 a and 
 by 73 b to be present in the solution of hydrogen sulphide 
 [R 374] ? Explain the actions in c and d in accordance with 
 the hypothesis of ions. 
 
 e. To 2-3 c.c. of potassium dichromate solution add dilute 
 sulphuric acid in large excess (?). What change may be 
 
74] SULPHUR 67 
 
 assumed to have taken place? Now saturate (test? Note 36, 
 below) the mixture with hydrogen sulphide (?). What are the 
 colors of the solution and of the precipitate, respectively? 
 Show that two kinds of ionic chemical change are here illus- 
 trated. 
 
 /. Take 2-3 c.c. of potassium permanganate solution and 
 treat exactly as in e (?). Answer the same questions (?). 
 What chemical property of hydrogen sulphide is illustrated in 
 e and / ? 
 
 Note 36. To learn whether a liquid is saturated with a gas, 
 withdraw the delivery tube while the gas is issuing, cover the 
 mouth of the vessel quickly with the thumb or the palm of the 
 hand in such a way that the gas above the liquid has no time to 
 escape (why?), and shake vigorously. If the thumb now adheres 
 to the mouth of the vessel, the liquid is not yet saturated (why?). 
 If the liquid is saturated, what will be the pressure of the gas over 
 it after shaking [R 153] ? If air is allowed to displace the gas over 
 the saturated liquid, what effect will be observed on shaking as 
 described above? 
 
 74. Properties of Aqueous Hydrogen Sulphide: II. Sul- 
 phides [Hood]. 
 
 . a. What was the reaction of aqueous hydrogen sulphide 
 towards litmus (73 a) ? What ionic substances are present? 
 
 b. Take about 6 c.c. of sodium hydroxide solution and 
 saturate (test? Note 36) it with hydrogen sulphide (?). Test 
 with litmus (?) and use in c, d, e, and /. How should you 
 proceed to prepare normal sodium sulphide (solid)? 
 
 c. To a few drops of the solution prepared in 6 add dilute 
 hydrochloric acid (?) (see 75 6). 
 
 d. To a few drops add some bromine-water (?). 
 
 e. To a larger portion of the same solution add a little pow- 
 dered roll sulphur and shake from time to time. Is sulphur 
 soluble in water? What is here to be inferred [Note 35, p. 48] ? 
 When the solution has become very yellow (?) in color, filter. 
 Acidify the filtrate with dilute hydrochloric acid (i.e., add 
 more than an equivalent amount of the acid) (?). Recall an 
 experiment with iodine which resembles this experiment with 
 sulphur (?). 
 
 /. Allow the remainder of the solution from b to remain 
 exposed to the air for several days (?). When a change in 
 color has occurred, add dilute hydrochloric acid in excess (?). 
 Explain. 
 
 g. Take five clean test-tubes and obtain 2-3 c.c. of the solu- 
 tion of each of the following substances. Dilute each specimen 
 with 10-20 c.c. of water and saturate (test?) with hydrogen 
 
68 SULPHUR [ 75 
 
 sulphide: (a) Cupric sulphate (?), (b) Lead nitrate (?), (c) Cad- 
 mium sulphate (?), (d) Zinc acetate (?), (e) Ferrous sulphate 
 (make a solution of ferrous-ammonium sulphate [R 754] and 
 use it for this [Note 37, below]). 
 
 Explain the precipitation in accordance with the hypothe- 
 sis of ions. 
 
 Pour away a part of the contents of each test-tube, add a 
 large excess of dilute hydrochloric acid, and shake (?). Explain 
 the results. Divide the metallic sulphides obtained in this 
 experiment into two classes, and characterize those classes. 
 
 What property of hydrogen sulphide has been illustrated in 
 a, b, and g ? 
 
 Note 37. Ferrous-ammonium sulphate [R 754] is preferred to 
 ferrous sulphate because it becomes less rapidly impure through 
 oxidation by the air. In making equations, disregard the ammo- 
 nium sulphate. 
 
 75. Ionic Chemical Changes: Formation of an Inactive 
 Acid. 
 
 a. To 2-3 c.c. of sodium acetate solution add an equivalent 
 amount of dilute sulphuric acid, and warm gently. One product 
 may be recognized by its odor (?). 
 
 Formulate, study and explain this action according to the 
 directions in 64 a (second par.). Write a simple equation ex- 
 pressing the chief change which occurs. 
 
 b. Consider the action between sodium-hydrogen sulphide 
 and hydrochloric acid in 74 c. Formulate, study and explain 
 this action as in 64 a (second par.), modifying the scheme of 
 formulation to suit the case. Does the escape of the hydrogen 
 sulphide assist materially in making the action more nearly 
 complete? Does it assist at all? What property does the 
 escape of the hydrogen sulphide as a gas show this substance 
 to possess? 
 
 c. Make a single general statement describing all the cases 
 in which, when ionogens are mixed, and no precipitation or 
 volatilization occurs, a fairly complete chemical change will 
 nevertheless take place (see 66). 
 
 Give a list of acids (Appendix VI) which might be formed 
 in accordance with the principle embodied in your statement. 
 
 Could any bases be formed in accordance with the same 
 principle? Illustrate (Appendix VI). Could any salts be so 
 formed (Appendix VI) ? 
 
 76. Hydrolysis. 
 
 a. Dissolve a single crystal of sodium sulphide in water, and 
 test the solution with litmus paper (?). What ionic substance 
 
77] SULPHUR 69 
 
 causes this reaction? Which of the two substances taken 
 is capable of furnishing this ion? Formulate the interaction 
 of the two original substances and explain it [R 375]. 
 
 6. Test with litmus paper the solution of sodium carbon- 
 ate (?). Explain (Appendix VI). Of what sorts of radicals 
 must a normal salt be composed in order that the solution may 
 show an alkaline reaction [R 344] ? 
 
 c. Test with litmus paper the solutions of cupric sulphate (?) 
 and ferric chloride (?), and account for the result [R 344]. Of 
 what sorts of radicals must a normal salt be composed in 
 order that the solution may show an acid reaction? Why is 
 sodium chloride solution neutral? Define hydrolysis. 
 
 77. Sulphur Dioxide. 
 
 a. Touch with a warm platinum wire a bit of sulphur, and 
 bring the wire with the adhering sulphur again into the flame. 
 Withdraw, and note the color of the flame of the burning 
 sulphur and also the odor (?). 
 
 b. Heat in a hard glass test-tube a few particles of iron 
 pyrites. What is the sublimate? What gas is evolved? 
 
 c. [Hood] If sulphur dioxide gas is not furnished in the 
 laboratory from cylinders of the liquid,* it may be prepared for 
 
 use in 78 or 83 as follows: Fit up a 
 
 flask (250 c.c.) as in Fig. 8 (p. 20), r N S, 
 
 and attach to it by rubber connections M>HL 
 
 a gas washing-bottle (Fig. 18). This SSf 
 
 bottle is to serve here as a safety J TL 
 
 vessel (why and how?) and is to be 
 left empty. If 78 is to be performed, 
 a second, similar gas washing-bottle 
 will be required to dry the gas, and 
 will contain enough concentrated sul- 
 phuric acid to immerse the longer tube 
 to a depth of half an inch. Attach ^ 
 
 the empty bottle with the shorter tube 
 next to the generating flask (why not the reverse?). Place 
 about 10 g. of copper nails in the flask. Test the apparatus 
 to see that it is air-tight. Add 10-15 c.c. of concentrated 
 sulphuric acid. 
 
 Heat the flask and contents by means of a sand bath. 
 Leave the cork out at first and suspend in the acid a thermom- 
 eter. Note the temperatures at which chemical action becomes 
 perceptible (?) and at which it is conspicuous (?). Why cannot 
 
 * If the gas is not so furnished, 77 and 78 should be taken up after 
 82. 
 
70 SULPHUK [I 78 
 
 dilute sulphuric acid be used? Connect the apparatus and 
 continue heating to obtain the gas needed in 78 and 83. 
 
 d. Allow the generating flask with its contents to remain 
 over night, and then examine and describe all the contents (?). 
 
 78. Molecular Weight of Sulphur Dioxide [Quant. Hood]. 
 Clean and dry a 250 c.c. flask and provide it with a tightly 
 fitting cork. Weigh the flask and cork. This gives the weight 
 of the flask filled with air. Now fill it completely with sulphur 
 dioxide, by downward displacement of air, cork, and weigh 
 again. To insure its being full, repeat this operation till no 
 increase in weight occurs. Finally, allow the gas to escape, 
 and determine its volume by filling the flask with water up 
 to the cork and weighing again. Observe the temperature 
 and pressure of the atmosphere. 
 
 To obtain the weight of the empty flask and its cork, subtract 
 from the weight of the vessel filled with air the weight, under 
 the observed conditions, of a volume of air equal to its content 
 (1 liter of pure, dry air weighs 1.293 g. under normal condi- 
 tions, or the G.M.V. holds 28.955 g. of air). 
 
 The difference between this corrected weight and that of 
 the flask filled with sulphur dioxide is the weight of the latter. 
 Reduce the volume of the gas to normal conditions and calcu- 
 late the weight of the G.M.V. (22.4 1.) and of 1 1. 
 
 Enumerate carefully all the sources of error to which you 
 should expect this way of determining the density of a gas to 
 be liable. In doing this, consider each detail of the operation 
 very critically. 
 
 79. Preparation of Sulphuric Acid (Two students working 
 together). Obtain a distilling-flask (25 c.c.), rubber connections 
 for a safety bottle, a screw-clamp, and a Chapman pump from 
 the storeroom. Fit up with your one-liter bottle the apparatus 
 as in Fig. 19. Charge the hard glass tube with about 10 g. of 
 granular pyrite,* and place a small, loose plug of- asbestos just 
 beyond the pyrite to retain any unburnt sulphur. Put into the 
 distilling-flask about 10 c.c. of pure concentrated nitric acid 
 [Side-shelf]. The safety bottle, half filled with water to show 
 the rate at which air is being drawn through the apparatus, is 
 attached to the water pump. The total air admitted is regu- 
 lated by the screw-clamp between the pump and the safety 
 bottle) while the proportions which pass over the pyrite and 
 
 * In case very pure pyrite is not available, it is better to depart 
 from the common manufacturing process by substituting a boat con- 
 taining a little sulphur. 
 
79] 
 
 SULPHUR 
 
 71 
 
 which carry over the nitric acid vapor, respectively, are regulated 
 by pinching one of the tubes with the finger and thumb. 
 
 First heat the pyrite in a very slow stream of air until the 
 sulphur burns. Then warm the nitric acid, and, by pinching 
 the tube admitting air to the pyrite-burner, divert part of the 
 air current so that it may carry over a little of the vapor of 
 the acid. Heat the pyrite strongly and continuously. Repeat 
 the introduction of air laden with nitric acid at intervals, when- 
 
 Fig. 19 
 
 ever the disappearance of the red fumes in the bottle shows 
 that a further supply is needed. 
 
 After a crust of white crystals (?) has formed in the bottle 
 (there may be considerable delay before crystallization starts), 
 remove the attachments and blow the gases from the interior 
 by means of the air blast. If crystallization fails to begin after 
 a reasonable time (note that an interaction even between the 
 molecules of gases may be slow, in spite of the completeness of 
 the mixing), the cause is either the introduction of too much 
 water along with the nitric acid, or the high temperature pro- 
 duced by the chemical actions taking place in the bottle. Re- 
 moving the attachments and cooling the bottle in a stream of 
 water frequently brings crystallization about. 
 
 Add 4-5 c.c. of water and wash down the crystals with it. 
 Describe all that' happens. If more of the product is required, 
 the apparatus may be connected up again and a further supply 
 of sulphur dioxide drawn into the bottle, and subsequently 
 
72 SULPHITE [ 80 
 
 more nitric acid vapor can be added. Finally any remaining 
 crystals may be decomposed by water. 
 
 Filter the liquid in the bottle through a very small filter 
 paper into a dish, rinsing the bottle with 2-3 c.c. of water, and 
 evaporate on the sand bath [Hood] until the liquid begins to 
 fume strongly (?). This will remove any nitric or nitrous acid 
 that it may contain. Use the result for 80 a and 6. 
 
 80. Properties of Sulphuric Acid. 
 
 a. Dip a match-stick into the liquid from 79, and make 
 marks on a sheet of paper. Set both paper and stick on a 
 radiator (?). What property of the acid is here observed? 
 
 b. After the acid prepared in 79 is cold, dilute it with 2-3 
 volumes of water. Test the solution with litmus paper (?). 
 Add to it barium chloride solution (?). To learn whether the 
 action is easily reversible, add pure hydrochloric acid to the 
 mixture (?). The formation of this precipitate, and its insolu- 
 bility, furnish a distinctive test for what radical? Which sub- 
 stances contain this radical? What property of sulphuric acid 
 is shown in 6? 
 
 c. Pass a current of hydrogen sulphide through 2-3 c.c. of 
 concentrated sulphuric acid. Note the odor (?) and precipi- 
 tate (?). What property of sulphuric acid is shown in c? What 
 other evidence of this property have we previously observed 
 (see 16 d and 52)? 
 
 d. Place a small piece of sulphur in 2-3 c.c. of concentrated 
 sulphuric acid, and heat (?). Note the odor (?). 
 
 e. Heat a little charcoal with sulphuric acid and note the 
 odor (?). What property of sulphuric acid is shown in d and 
 e? Why would not dilute sulphuric acid show this property? 
 
 /. [Hood] Take 2-3 c.c. of concentrated sulphuric acid in 
 a test-tube. Suspend a thermometer so that the bulb is com- 
 pletely immersed in the acid. Heat the contents of the tube 
 by means of a small flame and note the temperature at which 
 any effect (?) is observed, and that at which it is conspicuous. 
 Relate this temperature to that observed in 77 c. [CAUTION! 
 During the heating remember that, if the tube should crack, 
 the hot acid may splash on the clothes and hands and produce 
 severe burns. Exercise proper caution. Be careful not to 
 wash out this tube until the acid has cooled.] 
 
 81. Sulphuric Acid as a Dibasic Acid. 
 
 a. (Two students working together). Fill a burette with a 
 solution of potassium hydroxide made as in 66. Add 15 c.c. 
 of concentrated sulphuric acid slowly to 35 c.c. of water in a 
 beaker, and fill another burette with the diluted acid, when cold. 
 Ascertain as in 66 what volume of the alkali will neutralize 5 c.c. 
 
83] SULPHUR 73 
 
 of the acid. Concentrate the mixture by evaporating to about 
 10 c.c., remove the dish from the water bath, and allow the 
 resulting solution to crystallize. Dry the crystals on filter 
 paper. To a second portion of the acid (use no phenolphtha- 
 lem. Why?), twice as great as before, add exactly the same 
 amount of the alkali. Evaporate to about 5 c.c. and treat as 
 before. 
 
 Compare the two lots of crystals as regards (a) form, (b) 
 taste, (c) reaction of their solution towards litmus. Con- 
 firm by studying the same properties of the purer substances 
 found on the side-shelf. Explain the differences between your 
 own preparations and those on the side-shelf, if any are observed. 
 Explain, in terms of the hypothesis of ions, the difference in 
 the reactions of the two products towards litmus. 
 
 Formulate, study, and explain the actions, as in 64 a (second 
 par.). Take the components of the original solutions [R 346] 
 one by one, and describe what happens to each, (a) during 
 complete neutralization, (b) during semi-neutralization in the 
 second part of the experiment. 
 
 What would be the effect of heating perfectly dry specimens 
 of each of the salts [R 390]? 
 
 How many grams of potassium-hydrogen sulphate are re- 
 quired to give 1 liter of a solution of normal concentration 
 in respect to (a) potassium-ion, (b) hydrogen-ion, (c) sul- 
 phate-ion? 
 
 6. Make a solution of sodium bicarbonate and test it with 
 litmus paper (?) and with Congo red paper (?). Why do some 
 acid salts show little or no acid reaction towards indicators? 
 
 Define carefully, and illustrate, the terms: normal (neutral) 
 salt, acid salt, and basic salt [R 359]. 
 
 82. Sulphates. Place some ferric sulphate in a porcelain 
 crucible supported on the clay triangle, and heat strongly 
 [Hood] with the blast-lamp, continuing the heating after all 
 the water has been driven off(?). What are the properties of the 
 vapor given off (?), and what is it? What is the residue? Re- 
 late this result to that in 80 /. Recall action of heat on dehy- 
 drated gypsum (23 /). Classify the sulphates in accordance 
 with this distinction. 
 
 83. Properties of Sulphurous Acid [Hood]. Use a stream of 
 sulphur dioxide from a cylinder of the liquefied gas, or from the 
 apparatus described in 77 c. 
 
 a. Pass a stream of sulphur dioxide for a few minutes into 
 a test-tube full of water. Test the solution with litmus paper (?). " 
 What compound is present in the aqueous solution of the gas? 
 Note also the odor of the liquid (?). Formulate the interac- 
 
74 SULPHUR [ 84 
 
 tion, showing all the substances present. Divide the solution 
 into four portions. 
 
 b. Boil one portion persistently [Hood], noting from time to 
 time the odor and reaction towards litmus (?). In respect to 
 the result, does the solution of this gas resemble that of hydro- 
 gen sulphide, or that of hydrogen chloride? 
 
 c. To another portion add barium chloride solution (?). To 
 ascertain whether this action is easily reversible, add excess of 
 pure hydrochloric acid (?). 
 
 d. To the third portion add bromine-water (?) until the 
 color is permanent (?). Now add barium chloride to the mix- 
 ture (?) and then pure hydrochloric acid (?). 
 
 What chemical properties of sulphurous acid are illustrated 
 in a, 6, c, and d, respectively ? 
 
 e. Leave the fourth portion of the sulphurous acid for several 
 days exposed to the air in an open test-tube or bottle. Then 
 add barium chloride and pure hydrochloric acid, and compare 
 the result with that in c. 
 
 /. To 2-3 c.c. of potassium dichromate solution, add three 
 or four equivalents (a large excess) of dilute sulphuric acid. 
 What substances does the mixture contain? Lead a stream 
 of sulphur dioxide through the solution until no further change 
 is observed (?). 
 
 g. Take 2-3 c.c. of potassium permanganate solution, treat 
 as in / (?), and answer the same questions (?). 
 
 h. Fill a bottle with the gas by downward displacement, 
 introduce some moistened litmus paper or grass, and close 
 with a glass plate (?). 
 
 What properties of sulphurous acid are illustrated in e, /, g, 
 and h, respectively ? 
 
 84. Sulphites. 
 
 a. To 1 g. of sodium sulphite add any dilute mineral acid 
 (that is to say, one of the common acids, but not an organic 
 acid such as acetic) (?). Formulate this action completely. 
 
 6. Dissolve a minute amount of sodium sulphite in water 
 and add bromine-water in excess (test? The color must 
 remain). Remove the excess of bromine by boiling. Add 
 barium chloride solution (?) and then pure hydrochloric acid (?). 
 Compare this experiment with that in 83 d, consider whether it 
 is probably the molecular sodium sulphite, or the sulphite-ion, 
 which is oxidized, and make the equation accordingly. 
 
 c. Heat persistently about 1 g. of sodium sulphite in a porce- 
 lain crucible, supported on the clay triangle over a blast-lamp. 
 When cool, acidify with hydrochloric acid (?) and note the 
 odor. If any sulphur is precipitated, account for its formation. 
 
86] SULPHUR 75 
 
 d. Place about 1 g. of sodium bisulphite in a hard glass 
 test-tube and heat cautiously with the tube in a horizontal 
 position (why?). Note the odor (?) and whether any vapor 
 condenses (?). This behavior is typical of that of many acid 
 salts. Why is the behavior of potassium bisulphate somewhat 
 different (81 a)? What would be the final effect of heating 
 sodium bisulphite more strongly and for a longer time? 
 
 85. Thiosulphates. 
 
 a. Dissolve about 5 g. of sodium sulphite in about 20 c.c. of 
 water in a small flask. Add 4-5 g. of flowers of sulphur to the 
 solution and boil gently over a small flame for 10-15 minutes. 
 Filter off the clear yellow solution and divide into two parts. 
 
 b. To one portion add excess of any dilute mineral acid (?). 
 Note the odor. 
 
 c. Take 10 c.c. of potassium iodide solution and dissolve in 
 it a few small crystals of iodine [R 235] (?). Now pour the 
 second portion of the solution from a slowly into this mix- 
 ture [R 396] (?). How can this action be used for estimating 
 the quantity of free iodine, and what substance then acts as 
 the indicator? For what purpose was the potassium iodide 
 employed? 
 
 d. Heat persistently about 1 g. of sodium thiosulphate in a 
 porcelain crucible over a blast-lamp (?). Note the appearance 
 of the residue. When cold, add dilute hydrochloric acid (?) 
 and identify the products. If any odor was observed during 
 the heating, can you now account for it? 
 
 86. Reduction of Sulphur Compounds. Mix a pinch of any 
 salt of a sulphur acid with an equal amount of anhydrous 
 sodium carbonate. Slightly char the end of a match from 
 which the head has been removed, and rub the charred part, 
 which should be about an inch in length, with a heated crystal 
 of hydrated sodium carbonate. Moisten the above mixture 
 with water, place some of it on the end of the match, and heat 
 in the reducing region of a small Bunsen flame. Put the result 
 on a clean silver coin lying in a watch-glass and moisten with 
 one drop of water (?). Then add some dilute mineral acid and 
 note the odor (?). This, known as the "hepar" test, is a test 
 for sulphur in any form of combination. 
 
CHAPTER XII. 
 
 THE ATMOSPHERE, NITROGEN, AMMONIA. 
 
 87. Preparation of Nitrogen from Air. Place a large plug of 
 copper turnings in the center of a hard glass tube. Fit with 
 corks and short glass tubes, and connect with the short tube 
 of the aspirator (Fig. 7, p. 16). Fill completely with water 
 the bottle and long tube of the aspirator and close the latter 
 with a screw clamp [Storeroom]. Arrange a vessel to catch 
 the water discharged. Now heat the copper red-hot and then 
 partly open the clamp so as to allow the water to be syphoned 
 off in a slow stream. The air will pass over the heated copper. 
 What change does the copper undergo, and what collects in the 
 bottle? After three-fourths of the water has run out, close the 
 clamp, disconnect the hard glass tube, attach a delivery tube 
 in its place, elevate the end of the syphon, and replace the 
 nozzle by a small funnel supported by a clamp. Pour water 
 into the funnel, and drive the gas over into a bottle of water 
 inverted in the pneumatic trough. Ascertain whether this 
 gas supports combustion. Describe the gas as regards color 
 and smell. What other gaseous substances, besides nitrogen, 
 does this gas contain? 
 
 Other ways of removing oxygen from the air are described 
 elsewhere (14, 88, and 94 d). 
 
 88. Proportion (by Volume) of Oxygen in the Air [Quant.] 
 (from Cooley's Laboratory Studies). Provide a large test- 
 tube with a two-hole rubber stopper. Into one hole fit a short 
 piece of glass tubing terminating in a rounded nozzle, the tip 
 of which projects but little beyond the bottom of the stopper. 
 Into the other hole fit a short glass rod. Test the apparatus 
 for air-tightness. Connect the upper end of the glass tube with 
 a short-stemmed funnel by means of a piece of rubber tubing 
 (15 cm. long) as in Fig. 20, and support the funnel in a clamp. 
 
 Disconnect the test-tube, temporarily, and remove the glass 
 rod from the stopper. Prepare an alkaline solution of potas- 
 sium pyrogallate by mixing 3 c.c. of pyrogallic acid solution 
 with 20 c.c. of a solution of potassium hydroxide specially 
 prepared for this experiment [Side-shelf], and pour this into 
 the funnel. Now open the clamp and permit this solution to 
 fill the rubber and glass tubes completely down to the opening 
 of the nozzle. Replace the test-tube, fitting the stopper tightly 
 
 76 
 
88] ATMOSPHERE, NITROGEN, AMMONIA 11 
 
 into its mouth. Finally, reinsert the glass rod, and so inclose 
 a volume of air equal to the content of the test-tube and at the 
 temperature and pressure of the atmosphere. These opera- 
 tions should consume as little time as possible, as the alkaline 
 solution gradually absorbs oxygen from the air of the room, 
 and thereby becomes useless for further absorption. 
 
 Now open the clamp, taking care not to warm the test-tube 
 by handling. A few drops of the alkaline solution will enter 
 
 Fig. 20 
 
 the test-tube, and, as the oxygen is absorbed, more of the 
 solution will flow in. Close the clamp and invert the test-tube 
 once or twice in order to bring the liquid thoroughly in con- 
 tact with the inclosed air. Finally, while the test-tube is in 
 the inverted position, reopen the clamp and equalize the 
 levels of the liquid in test-tube and funnel by raising or lower- 
 ing the former. Then close the clamp, restore the test-tube 
 to its original position, and mark the positions of the surface 
 of the liquid and of the bottom of the stopper by means of 
 paper labels or rings cut from rubber tubing. 
 
78 ATMOSPHEKE, NITEOGEN, AMMONIA [ 89 
 
 Disconnect the test-tube and wash out the liquid, taking care 
 not to get the alkaline solution upon the hands. Then, by 
 means of a burette filled with water, measure the volumes 
 required to fill the test-tube up to the lower and upper marks 
 respectively. The former is the volume of the oxygen, the 
 latter that of the air. Calculate the percentage of oxygen in 
 the air by volume. 
 
 89. Other Components and Density of Air. 
 
 a. Place 2-3 c.c. of clear barium hydroxide solution in the 
 bottom of a small beaker and leave it exposed to the air for 
 some hours (?). 
 
 b. Blow air from the lungs through a glass tube into 2-4 c.c. 
 of clear barium hydroxide solution (?). Explain. 
 
 c. How may the presence of aqueous vapor in air be demon- 
 strated (27 a)? How may its quantity be measured? 
 
 d. The weight of a measured volume of air, or of nitrogen, 
 may be determined by the method described in 112. From 
 the data so obtained the density may be calculated. 
 
 90. Nitrogen. 
 
 a. Place about 10 g. of pure sodium nitrite and about 8 g. 
 of ammonium chloride in a 250 c.c. flask fitted with safety and 
 delivery tubes (Fig. 8, p. 20). If b is to be performed, insert 
 between the flask and the delivery tube a U-tube containing 
 calcium chloride for drying the gas. . 
 
 Clamp the flask by the neck to a ring-stand, add about 15 c.c. 
 of water, and warm gently [CAUTION!]. As soon as the action 
 begins remove the flame, bring a dish of cold water under the 
 flask, and cool it for a few seconds at a time so that the action 
 may not become too violent, but may run uniformly. After 
 sufficient time has been allowed for the displacement of air 
 from the apparatus, fill a bottle with the gas over water in a 
 pneumatic trough. Has the gas odor or color? Does it sup- 
 port combustion? 
 
 b. (Two students working together). In a piece of hard 
 glass tubing, fitted with corks and short pieces of glass tubing, 
 place a porcelain boat half filled with powdered magnesium. 
 Connect one end of this tube to the outlet tube of the flask in 
 a generating nitrogen. Heat the magnesium strongly (two 
 Bunsen burners may be necessary). What action is noticed? 
 What is the color of the product? Transfer the contents, when 
 cold, to a dry test-tube, close with a cork, and use in 91 b. 
 May the same product be formed when magnesium is heated in 
 the air? What other compound would be formed under the 
 latter conditions? 
 
92] ATMOSPHERE, NITKOGEN, AMMONIA 79 
 
 91. Ammonia. 
 
 a. Fit a small flask with a cork and L-shaped exit tube, and 
 fonnect the latter with a U-tube. Put a little water in the 
 bend of the latter so as just to close the passage. Test the 
 apparatus for air-tightness. Place in the flask a mixture of 
 powdered quicklime and ammonium chloride, about 10 g. of 
 each, and warm gently [Hood]. After the air has been dis- 
 placed, the whole of the gas should dissolve in the water. Use 
 this solution of ammonium hydroxide in 92. 
 
 If required to determine the density of ammonia, which of 
 the various methods, should you select (15 a, 78, or 112)? 
 
 b. To the product from 90 b add a little water, boil, and 
 note the odor (?). What other substances would behave in the 
 same way as magnesium nitride? 
 
 e. Heat a small piece of gelatine in a dry test-tube and note 
 the odor (?). 
 
 92. Ammonium Hydroxide: An Inactive Base: Salts of 
 Amm o ilium . 
 
 a. Test the reaction towards litmus paper of the solution 
 made in 91 a (?). What substance is present? Hold a glass 
 rod dipped in concentrated hydrochloric acid over the solu- 
 tion (?). What substance is shown by this test to be present? 
 Formulate the relations (in equilibrium) of all the substances 
 in the solution. 
 
 b. Expose 1-2 c.c. of the solution from 91 a in an evaporat- 
 ing-dish for 24 hours and then note the odor (?) and test with 
 litmus (?). Heat [Hood] another small portion in an evapo- 
 rating-dish for some minutes, and note the odor from time to 
 time (?). Does the aqueous solution of this gas behave like 
 that of hydrochloric acid, or like that of hydrogen sulphide? 
 What other gases resemble ammonia in this respect? 
 
 c. Neutralize (test ?) the remainder of the solution from 91 a 
 with dilute sulphuric acid and evaporate to dryness on a water 
 bath (?). Formulate this action after the model in 64 a (sec- 
 ond par.), and explain what happens during neutralization to 
 each of the substances present in the ammonium hydroxide 
 solution. 
 
 Scrape the residue into the middle of the dish, invert over it a 
 small funnel, the stem of which has been closed with a paper 
 plug, and heat the dish strongly and persistently (?). To learn 
 whether the sublimate is identical with the residue, examine it 
 for the ammonium radical (92 d, below), and the sulphate (80 6) 
 radical. 
 
 d. Dilute 1-2 c.c. of ammonium chloride solution and add to 
 ft 1-2 c.c. of sodium hydroxide solution (?). Note the odor (?). 
 
80 ATMOSPHERE, NITROGEN, AMMONIA [ 92 
 
 Formulate, study, and explain this action according to the 
 directions in 64 a (second par.). Include in the formulation the 
 liberation of the ammonia. What is an inactive base? 
 
 The evolution of ammonia when an active base is added to a 
 solution of any ammonium salt is used as a test for the latter. 
 Why would not simply heating the salt by itself (as in 92 c and e) 
 serve as a test? What salt of ammonium that we have used 
 gives, when heated, no ammonia? 
 
 e. Place some ammonium chloride in the middle of an open, 
 hard glass tube. Support this in a very slightly inclined (5) 
 position by means of a clamp attached close to one end. Put 
 pieces of moistened litmus paper (both colors) in each end, and 
 heat the salt strongly (?). Watch closely the effect on the 
 papers at each end (?). What does this experiment show to be 
 the action of heat on ammonium chloride? Which gas appeared 
 first at the ends of the tube, and why first at both lower and 
 higher ends? What should be the relative speeds, by calcula- 
 tion, of the two gases [R 108] ? Why does the second gas reach 
 the ends so very much later than one might expect? Has gravity 
 any influence on the result? 
 
CHAPTER XIII. 
 
 OXIDES AND OXYGEN ACIDS OF NITROGEN. 
 
 93. Preparation of Nitric Acid [Hood]. Pulverize about 10 g. 
 of sodium nitrate and place it in a dry retort or distilling-flask 
 [Storeroom]. Cover the salt with concentrated sulphuric acid 
 and wait until this liquid has permeated and moistened the 
 entire mass. Support the vessel by a clamp upon a sand bath 
 and allow the neck of the retort (or side-tube of the flask) to ex- 
 tend to the bottom of a small flask (the receiver). In order that 
 the vapor of the nitric acid may condense, this flask is immersed 
 in a vessel of cold water, and is covered with filter paper which is 
 continually moistened. Heat the retort gently until a sufficient 
 amount of nitric acid has been obtained, and preserve the prod- 
 uct in a corked test-tube for use in 97 a [CAUTION! Serious 
 wounds may be caused by allowing this liquid to come in contact 
 with the skin]. What is the composition of "concentrated" 
 nitric acid [R 440], and how does the acid here made differ from 
 the " concentrated " acid ? What law teaches us to expect that 
 the former will be more active (and dangerous to handle) than 
 the latter? 
 
 94. Nitric Oxide. 
 
 a. Into a small flask, fitted as in Fig. 8 (p. 20) and contain- 
 ing about 10 g. of copper nails, pour 10-15 c.c. of water 
 and then an equal amount of concentrated nitric acid [Desk]. 
 Collect the gas evolved over water in the pneumatic trough. 
 Note the color of the gases first formed, and observe the color 
 of the pure gas finally collected over the water. Fill two 
 bottles (for 6, c) and the large test-tube (for d) with the gas. 
 
 The production of colored gases, when copper is added, is 
 used as a test for nitric acid. 
 
 6. Into one bottle introduce a burning match (?). Note 
 also the effect of opening this bottle to the air (?). This effect 
 may be used as a test for nitric oxide, or, conversely, for 
 oxygen. -How should you apply it for each of these pur- 
 poses? 
 
 c. Into the second bottle lower a deflagrating spoon with a 
 very little burning, red phosphorus (?). What change in 
 volume does the gas undergo in this action? 
 
 d. Transfer without loss the nitric oxide in the test-tube to 
 a small beaker filled with water and inverted in the pneumatic 
 
 81 
 
82 ACIDS OF NITKOGEN [ 95 
 
 trough. Fill the test-tube with oxygen and add this to the 
 nitric oxide (?), thus securing equal volumes of the gases. 
 Shake the beaker, keeping the mouth well immersed, until no 
 further change occurs. Transfer the remaining gas, which 
 must be either nitric oxide or oxygen, back to the test-tube. 
 Note its volume, and determine what gas it is (?). What 
 relative volumes are (a) used for the interaction and (b) 
 required by the equation? 
 
 How might the proportion of oxygen in the air be deter- 
 mined by application of this interaction? 
 
 e. Prepare 10 c.c. of a concentrated solution of ferrous- 
 ammonium sulphate [Note 37, p. 68] (or ferrous sulphate) and 
 divide into three portions. Into one pass a gentle stream of 
 nitric oxide (?). Boil the liquid (?). 
 
 /. Add to the second portion of ferrous-ammonium sulphate 
 solution an equivalent amount of dilute sulphuric acid, heat 
 to boiling, and add nitric acid [Desk] drop by drop (mix after 
 each drop) until there is no further action (?). What gas is 
 liberated? 
 
 g. Dissolve a single crystal of sodium nitrate in 1-2 c.c. of 
 water, and add a part of this solution to the remainder of the 
 ferrous-ammonium sulphate solution from e. Pour concen- 
 trated sulphuric acid cautiously and steadily down the side of 
 the test-tube until it forms a considerable layer at the bottom 
 of the tube. Note the brown ring between the layers, and 
 explain by reference to e and /. This constitutes a delicate 
 test for nitric acid or a nitrate. In what way does the avoid- 
 ance of mixing (and consequent ring-formation) contribute to 
 the delicateness of the test? 
 
 h. Take 2-3 c.c. of concentrated nitric acid in a test-tube, 
 warm it slightly, and lead through it a stream of nitric 
 oxide (?). By waiting until the air has been wholly displaced, 
 make certain whether the colored gas is formed by an inter- 
 action of the nitric oxide with the nitric acid, or simply by 
 contact with the air in the test-tube (?). 
 
 95. Nitrogen Tetroxide. Fit a hard glass test-tube with a 
 cork and L-shaped exit tube, and see that it is air-tight. Place 
 8-10 g. of lead nitrate in the tube and clamp it in a horizontal 
 position. Prepare a strong solution of sodium hydroxide by 
 dissolving 2 g. of the solid in 7 c.c. of water and allow the 
 delivery tube to reach the bottom of this solution. Now heat 
 the lead nitrate (?) persistently until gas is no longer given off, 
 or until the liquid is no longer soapy to the touch. If all the 
 gas is not absorbed in the alkali, test the escaping bubbles for 
 oxygen (?). Use the solution in 100 d. 
 
97] ACIDS OF NITROGEN 83 
 
 What is the residue? What other nitrate have we decom- 
 posed in the same way (36)? This behavior when heated is 
 typical of the nitrates of the heavy metals, and may be used 
 for identifying them. Which nitrates, when heated, would 
 leave the metal, instead of the oxide [R 362] ? Why do com- 
 mercial specimens of mercuric oxide [R 657], when heated in 
 the preparation of oxygen, frequently give off a brown gas? 
 
 96. Principles Involved in Making Nitric Acid. 
 
 a. Addition of copper (94 a) to a liquid containing nitric 
 acid gives red vapors. Is this a test for any substance con- 
 taining the nitrate radical, or only for nitric acid? Dissolve a 
 little sodium nitrate in water, add a copper nail, and warm (?). 
 
 b. Was nitric acid formed in 93 on mixing sodium nitrate 
 and sulphuric acid, before the distillation began? Solve this 
 question by mixing sodium nitrate (finely pulverized) with 
 concentrated sulphuric acid, adding a very little water [CAUTION] 
 and agitating for a minute or so. Lower a glass rod dipped in 
 ammonium hydroxide solution into the test-tube (?). Apply 
 also the copper test as in a (?). 
 
 c. Is the action of sulphuric acid upon sodium nitrate rever- 
 sible? Take 2-3 c.c. of. concentrated sodium-hydrogen sulphate 
 solution and add to it an equal or even greater volume of pure 
 [Side-shelf] concentrated nitric acid. Cool the mixture in a 
 stream of cold water and stir with a glass rod (?). The salt 
 solution must be sufficiently concentrated, otherwise the ex- 
 periment will fail. Examine with a lens the crystalline product 
 that separates out (?). If the action is reversible, what 
 enabled us to obtain a large yield of nitric acid in 93? 
 
 d. Will other acids behave like sulphuric acid? Mix some 
 pulverized sodium nitrate with phosphoric acid, agitate for a 
 minute or two, and apply the copper test as in a (?). Could 
 phosphoric acid be used in the preparation of nitric acid as in 
 93? Why? Could hydrochloric acid be used [R 182, 440, 
 448]? Could hydrofluoric acid be used [R 241, 440]? Give 
 reasons for your answers. 
 
 97. Properties of Nitric Acid. 
 
 a. If 93 was performed, note the color of the sample (?). 
 Blow some air from the blast through the acid (?). What is 
 the color of pure nitric acid? To what substance was the color 
 of this sample due, and by what action was it formed? What 
 property of nitric acid does its formation indicate? Blow the 
 moist air of the breath over this specimen (?). What property 
 does this indicate? Determine the boiling-point of anhydrous 
 nitric acid by boiling this sample in a distilling-flask with a 
 thermometer immersed in the vapor (?). 
 
84 ACIDS OF NITROGEN [ 97 
 
 6. Test dilute nitric acid with litmus paper (?). 
 
 c. Add 1 g. of sulphur to 2-3 c.c. of concentrated nitric acid 
 and boil [Hood] for two or three minutes (?). Is there any 
 evidence of action? Pour the clear liquid into another test- 
 tube, dilute with water, and test for the sulphate radical (?). 
 What property of nitric acid is here shown? 
 
 d. The interaction of dilute (16 c) and of concentrated 
 (16 d) sulphuric acid with metals has already been studied. 
 Try the action of (a) magnesium and (b) zinc, separately, 
 either upon dilute or upon concentrated nitric acid, and of 
 (c) copper upon both, and explore the whole action thoroughly 
 in each of the four cases as follows : 
 
 Fit up a side-neck test-tube with a dropping-funnel passing 
 through a perforated cork, and attach a delivery tube (or use 
 an ordinary test-tube with cork and delivery tube). Place a 
 small amount of the metal in the tube and admit the acid 
 from the dropping-funnel. After the air has been displaced, 
 collect the gas over water in a test-tube inverted in the pneu- 
 matic trough (or in an evaporating-dish). 
 
 If the gases in the generating tube remain colored, nitrogen 
 tetroxide is present. If the gas is at first colored, but becomes 
 colorless, nitric oxide is present: confirm by admitting air to 
 the tube in which it has been collected (?). Hydrogen, if 
 liberated alone, would show neither of these properties, but 
 might be identified by its inflammability (?). If both hydro- 
 gen and nitric oxide are liberated, why may we not feel confi- 
 dent of being able to cause the former to burn when thus 
 diluted? How can the nitric oxide be separated (94 d or e) 
 so as to leave the hydrogen in an inflammable condition? 
 Examine the gas in each case for these three substances. 
 
 If ammonia is formed in any of these four cases by complete 
 reduction of the nitric acid, where will it be found, and in what 
 condition? (Consider all the circumstances carefully, or you 
 will answer wrongly.) Test for its presence (92 d) (?). Evap- 
 orate on a water bath [Hood] the solution remaining in the test- 
 tube after any one of these experiments (?). Collect the residue 
 and heat it in a dry test-tube (?). In what form of combination 
 was the metal (95)? 
 
 e. Grasp a test-tube by means of a strip of folded paper (Fig. 5, 
 p. 14), place in it a piece of granulated tin, and add some con- 
 centrated nitric acid(?). When the action has exhausted itself, 
 add much water and boil. Fit a small filter paper carefully into 
 a funnel [Note 24, p. 9], collect the solid product upon the filter 
 and wash [Note 38, below] precipitate and paper with water. 
 Test the filtrate as it runs through to see whether it becomes 
 
99] ACIDS OF NITKOGEN 85 
 
 neutral (?). Spread the filter paper with the precipitate upon a 
 water bath or radiator to dry and ascertain as in 95 whether the 
 product is a nitrate (?). Why would the test in 94 g be unsuit- 
 able here? 
 
 /. Comparing dilute and concentrated nitric acid with the 
 same forms of sulphuric acid, what points of resemblance and of 
 difference have you observed? 
 
 Account for the differences in the actions of dilute and of con- 
 centrated nitric acid in c, d, and e, in respect to the gaseous prod- 
 uct of reduction which is typical of each. 
 
 g. Dip a piece of wool in concentrated nitric acid, or note 
 the nature of the products formed when the acid falls upon 
 the hands or clothing (?). Would cotton behave in the same 
 way? Explain [R 441]. 
 
 h. To 1-2 c.c. of nitric acid add 3-4 c.c. of pure, concen- 
 trated hydrochloric acid, warm gently, and note the appear- 
 ance and odor (?). Will these acids, singly, attack all metals? 
 Explain (?). What metals are attacked only by the mixture 
 (aqua regia), and why? What form of combination do the 
 metals assume? 
 
 Note 38. In filtering, observe the directions in Note 24. To 
 wash a precipitate first let the mother liquor drain away completely 
 and press the precipitate with a spatula. Then cover the contents 
 of the funnel, including the whole paper (why?), completely (and 
 repeatedly, if necessary) with the washing material. When the 
 washing is complete, dry the product by pressing with dry filter 
 paper. 
 
 98. Nitrous Oxide. 
 
 a. In a test-tube or small flask provided with a cork and 
 delivery tube (Fig. 10, p. 24), place 10 g. of ammonium nitrate. 
 Heat cautiously with a small flame [CARE!], and, after the air 
 has been expelled, collect the gas in two bottles over water 
 (warm water, if available. Why?). Why is the stream of 
 gas so slow, relatively to the vigor of the action? What are 
 the relative volumes of the products at 100 ? 
 
 b. Into one bottle lower a glowing splinter of wood (?). 
 
 c. Into the other bottle lower a deflagrating spoon contain- 
 ing a very little burning, red phosphorus (?). If this experi- 
 ment were to be conducted in a closed vessel, what, if any, 
 change in pressure would be observed after the products had 
 cooled? 
 
 99. Nitrates. 
 
 a. How does lead nitrate behave when heated? What other 
 nitrates behave like it? How does ammonium nitrate behave 
 when heated? 
 
86 ACIDS OF NITROGEN [ 100 
 
 b. Take 2-3 g. of sodium nitrate in a hard glass test-tube, 
 and heat strongly and persistently until the evolution of gas 
 ceases. Test the escaping gas for oxygen. Use the residue in 
 100 a. What nitrates behave in this way when heated? 
 
 100. Nitrites and Nitrous Acid. 
 
 a. When the residue from 99 b has cooled, add not more than 
 3 c.c. of water, shake vigorously until the whole has dissolved, 
 and divide into three parts. To one portion add dilute sul- 
 phuric acid (?). How could a nitrite be distinguished from a 
 nitrate? 
 
 b. To 5 c.c. of starch emulsion add a drop of potassium iodide 
 solution and some dilute sulphuric acid, and then introduce a 
 little of the solution from a (?). 
 
 c. Dilute 1-2 c.c. potassium permanganate solution with a 
 large excess of dilute sulphuric acid, and add a drop of this 
 mixture to the solution from a (?). The actions in a, 6, and c 
 serve for the detection of nitrites and nitrous acid. What other 
 characteristic property of nitrites has been encountered (90 a) ? 
 
 d. Examine now the solution obtained by leading nitrogen 
 tetroxide into sodium hydroxide (96). Take a portion of the 
 solution and acidify (test?) with dilute sulphuric acid (?). To 
 the mixture add a drop of diluted potassium permanganate 
 solution (?) or some starch emulsion containing a drop of potas- 
 sium iodide (?). What substance was present? 
 
 To ascertain whether this alkaline solution contains a nitrate 
 as well as a nitrite, the latter must first be eliminated. To the 
 remainder of the alkaline solution, transferred to a small flask, 
 add at least 5 g. of ammonium chloride [Hood] and heat to 
 boiling (?). Explain the evolution of ammonia. What other 
 gas is given off (90 a) ? When the action is entirely over, add 
 about 10 c.c. of water, and shake. Prepare 2-3 c.c. of a concen- 
 trated solution of ferrous-ammonium sulphate [Note 37, p. 68] 
 (or of ferrous sulphate), add to it 1-2 c.c. of the solution to be 
 examined, and complete the test for a nitrate described in 
 94 g (?). Write now the equation for the action of nitrogen 
 tetroxide on sodium hydroxide. 
 
 101. Active (Nascent) State of Hydrogen. Dilute some 
 potassium permanganate solution with water, and add an equal 
 volume of dilute sulphuric acid to it. Divide into two parts. 
 Through one pass a stream of hydrogen gas from the laboratory 
 supply or from a Kipp's apparatus (?). To the second add some 
 zinc dust, and shake (?). Interpret the result. 
 
CHAPTER XIV. 
 
 PHOSPHORUS. 
 
 102. Phosphorus. What difference in general chemical 
 behavior do the two allotropic modifications of phosphorus 
 exhibit [R 459] ? What product is formed when phosphorus 
 burns in excess of oxygen, and what is formed when the oxygen 
 is limited in amount [R 464] ? 
 
 103. Phosphine [Hood]. Place a very small piece of cal- 
 cium phosphide in a little water in a beaker (?). What mode 
 of forming ammonia is similar to this, and how is it similar? 
 Repeat, using dilute hydrochloric acid instead of water (?). 
 What mode of forming hydrogen sulphide is similar to this, 
 and how is it similar? In what ways does phosphine differ 
 from ammonia [R 461] ? Which of the differences are shown 
 in these experiments? Why should we expect ammonia and 
 phosphine to be alike? Test with litmus the water in which 
 the calcium phosphide was placed (?), and explain. 
 
 104. Metaphosphoric Acid. 
 
 a. Throw 2 g. of phosphorus pentoxide, in minute portions 
 at a time, into a beaker containing 10 c.c. of cold distilled 
 water (?). Allow the liquid to stand for a few minutes, or 
 until clear. Test the liquid with litmus paper (?). What 
 acid is present [R 464] ? Use the solution for- b and c and 105 a. 
 Why could not the product made by burning phosphorus in a 
 closed volume of air be used in this case? 
 
 6. To a part of the solution from a add silver nitrate solu- 
 tion, a little at a time, shaking between additions, until a per- 
 manent precipitate (color?) is formed (?). Is an action like 
 this reversible? Is this action incomplete? To answer, add 
 one or two drops of diluted (1 : 4 Aq.) ammonium hydroxide 
 solution, observe whether the precipitate increases, and ex- 
 plain. This is a reaction of what ion? What other substances 
 would give it? 
 
 c. To 1 c.c. of albumen solution, add one or two drops of 
 the solution from a [R 468] (?). This is a reaction of the free 
 acid only, and not of metaphosphate-ion. 
 
 105. Orthophospnoric Acid. 
 
 a. Dilute the remainder of the solution of metaphosphoric 
 acid from 104 a with 10 c.c. of water, and boil in a small flask 
 vigorously for an hour or more. If necessary add more water 
 
 87 
 
88 PHOSPHORUS [ 106 
 
 to make up for the loss by evaporation. Cool the solution, 
 treat portions of it with silver nitrate and with albumen, as in 
 104 6 and c, and answer the same questions. How may the 
 formation of this acid, during the boiling, be hastened [R 465]? 
 
 6. Heat 1 g. of red phosphorus with 5 c.c. of slightly diluted 
 nitric acid in a test-tube (?). When the action has ceased, 
 filter, if necessary, and drive off the water and excess of nitric 
 acid completely by evaporation [Hood] on a water bath. Mix 
 with a few drops of concentrated nitric acid the syrup which 
 remains, and evaporate once more. It is essential that all the 
 nitric acid be finally removed (why?). Redissolve the syrup 
 in water and test with litmus paper (?). 
 
 Treat part of the solution with silver nitrate as in 104 b, 
 and answer the same questions (a black precipitate with silver 
 nitrate is due to phosphorous acid [R 469] formed by incom- 
 plete oxidation). 
 
 c. Take 5 c.c. of the solution of ammonium molybdate in 
 nitric acid [Side-shelf], add to it two drops of the solution of 
 orthophosphoric acid prepared in b, and warm gently (?). 
 This is a very delicate test for phosphoric acid, or orthophos^ 
 phate-ion. 
 
 106. Phosphates. 
 
 a. Take some sodium phosphate (secondary sodium ortho- 
 phosphate) solution and test it with litmus paper (?). What 
 ions are present? Are acid salts always acid towards litmus 
 (74 6)? If not, explain why they are not. Explain also the 
 actual reaction of this solution. 
 
 b. To a small part of the sodium phosphate solution add 
 silver nitrate solution until the precipitation is complete (?). 
 What is the precipitate (see 106 a and b) ? What are the reac- 
 tions towards litmus of the sodium phosphate (a) , and silver 
 nitrate solutions singly? Test the mixture with litmus 
 paper (?). What ion is evidently formed, and what product 
 is therefore present? 
 
 c. Prepare a little " magnesia mixture" by adding to 1 c.c. 
 of magnesium sulphate solution a few drops of ammonium 
 hydroxide and then excess of ammonium chloride solution. 
 Add to the rest of the sodium phosphate solution a little of 
 this mixture (?). In making the equation, disregard the 
 ammonium chloride, which is added only to prevent precipita- 
 tion of magnesium hydroxide. This is another test for the 
 presence of what ions? 
 
 d. Strongly heat [Blast-lamp] 2 g. of dry sodium phosphate 
 in an open crucible for twenty minutes, or until no further 
 change is observed (?). Dissolve the cold mass in water (it 
 
107] PHOSPHORUS 89 
 
 dissolves very slowly), and test a portion of the solution with 
 silver nitrate solution (?). This is a reaction of what ion? 
 Take the remainder of the solution, liberate the acid (?) by 
 adding acetic acid, and introduce a few drops of the liquid 
 into a solution of albumen (?). 
 
 Make a table showing the effects of the three phosphoric 
 acids upon albumen and the colors of their silver salts. How 
 could you identify salts of the three acids? 
 
 e. Heat strongly [Blast-lamp] 2 g. of microcpsmic salt as 
 in d. Note the odor (?) and reaction towards moistened litmus 
 paper (?) of the vapors given off. Dissolve the residue in cold 
 water and use the tests tabulated in d to learn what salt has 
 been formed. 
 
 /. Make a bead on a straight platinum wire as in 3 c, using 
 microcosmic salt instead of borax. Of what must the bead 
 consist? Now fuse with it (3 d) a minute particle of cupric 
 oxide (?). What difference do your experiments show as to 
 the relative stabilities of NaP0 3 , NaNO 3 (99 6), and KC10, 
 (16 a)? 
 
 107. Halides of Phosphorus. 
 
 a. What were the actions of water on the tribromide (44 6) 
 and tri-iodide (48 6) of phosphorus? 
 
 b. Place 2-3 c.c. of phosphorus trichloride in a dry test- 
 tube and blow the breath over the tube (?). Add water drop 
 by drop (?) until about 5 c.c. has been used, then boil the solu- 
 tion (Fig. 5, p. 14). Test the vapor with litmus paper (?) 
 and with a rod dipped in ammonium hydroxide (?). Evapo- 
 rate [Hood] the solution to a syrup on the water bath (?). 
 
 c. Transfer part of the syrup from b to a small test-tube and 
 heat until the gas evolved may be ignited at the mouth of the 
 tube (?). Note the odor of this gas before igniting it (?). 
 Continue heating until the action is over. Ascertain according 
 to 106 d which of the phosphoric acids constitutes the 
 residue (?). 
 
 d. Dilute 1 c.c. of potassium iodide solution with water and 
 dissolve in the liquid a crystal of iodine. Add to this solution 
 a little of the syrup from b, and shake (?). 
 
 e. Place upon a watch-glass 5 or 6 small granules of phos- 
 phorus pentachloride and blow the breath over them (?). 
 Throw them into 2-3 c.c. of water in a test-tube (?) and boil. 
 Test the solution with litmus (?). To a small part of the cold 
 solution add excess (test?) of silver nitrate solution (?). Fil- 
 ter. What remains upon the paper (?). To the filtrate add 
 diluted (1: 4 Aq.) ammonium hydroxide drop by drop (shake 
 between drops) (?). 
 
CHAPTER XV. 
 
 CARBON. 
 
 108. Charcoal. 
 
 a. Place a small piece of charcoal in a test-tube half full of 
 water (?). Now hold it under the water with the reverse 
 end of the file, and boil the water for several minutes (?). 
 When the whole has cooled, test once more the tendency of 
 the charcoal to float (?). Explain. 
 
 b. Boil dilute solutions of litmus and indigo, separately, 
 with pulverized animal charcoal, and filter each liquid (?). 
 The activity of the charcoal is much increased by previous 
 heating in a covered crucible. 
 
 c. Fit a hard glass test-tube with a cork and L-tube. Mix 
 Intimately in the mortar 1 g. of pulverized cuj)ric oxide with 
 1 g. of pulverized wood charcoal and place in tne tube. Heat 
 persistently [Blast-lamp] and pass the gases through 5 c.c. of lime- 
 Water (?). Examine the residue by rubbing vigorously under 
 water in the mortar and washing away the lighter particles (?). 
 What property of carbon is here illustrated? 
 
 109. Carbon Dioxide. 
 
 a. Place a few small pieces of magnesite (magnesium car- 
 bonate) in a hard glass test-tube fitted with a cork and L-tube 
 and heat strongly. Pass the gas through 5 c.c. of lime-water 
 in a test-tube (?). What is the residue [R 643] ? 
 
 6. Fit up a generating-flask (Fig. 8, p. 20) and connect with 
 two wash-bottles (Fig. 18, p. 69) containing water and con- 
 centrated sulphuric acid, respectively (what is the use of each 
 of these? The latter is unnecessary if 110 is omitted). Place 
 in the flask some pieces of marble and pour upon them diluted 
 hydrochloric acid. Collect the gas in two bottles by down- 
 ward displacement. 
 
 What substances could be substituted for marble in this 
 experiment? What other sources of carbon dioxide have been 
 encountered (13 c, 108 c, 109 a; see also 111 a and c, 113 b, 
 114 b, 118)? 
 
 c. To one add a little water, close with the hand, and 
 shake (?). Is the gas soluble in water, or not ? 
 
 Use the second to compare its weight with that of air. 
 Employ baryta-water as a test. 
 
 d. Take two test-tubes containing distilled water, pass a 
 
 90 
 
112] CARBON 91 
 
 current of the gas into one, and test each with litmus paper (?) . 
 Boil the solution, and test again with litmus (?). Explain. 
 
 e. Lead the gas into a little sodium hydroxide solution in a 
 test-tube until the solution is saturated (test? Note 36, p. 67) . 
 Let the solution stand until it dries spontaneously (first residue) . 
 Heat the dry residue (?) in a test-tube, and determine what two 
 things are given off. 
 
 To this residue after heating (second residue) add dilute 
 hydrochloric acid until all action (?) ceases. Evaporate the 
 solution on the water bath, and examine and taste this final 
 residue (?). Having recognized the products of the last action, 
 and taking into account the preceding observations, state what 
 the nature of the second and first residues must have been. 
 Write equations for all actions. 
 
 110. Molecular Weight of Carbon Dioxide [Quant.]. 
 
 a. Determine the weight of a measured volume of the gas by 
 the method used for sulphur dioxide (78), and calculate the 
 weight of the gram-molecular volume. What further informa- 
 tion must we have to enable us to determine the formula? 
 
 b. Use the quantitative results obtained in the synthesis of 
 carbon dioxide (34) along with this molecular weight, to calcu- 
 late the weights of carbon and of oxygen in a molecular weight 
 of the gas. What further steps are necessary in order to fix 
 the atomic weight of carbon? 
 
 111. Carbon Monoxide. 
 
 a. Heat about 10 g. of oxalic acid crystals with concen- 
 trated sulphuric acid in a generating-flask (Fig. 8, p. 20), and 
 fill a bottle over water with the gas which is given off. Shake 
 with lime-water (?). With what substance should we wash 
 the gas to remove the carbon dioxide? Arrange a wash-bottle 
 to purify the gas. Fill two bottles with the purified gas over 
 water. Test one with lime-water again (?). If the gas is 
 pure, burn that in the other bottle, add lime-water at once, 
 close quickly, and shake (?). 
 
 b. Devise a way of ascertaining roughly the relative volumes 
 of the two gases generated in a, and measure the proportion in 
 a test-tube full of the mixed gases. 
 
 c. Pass a stream of the purified carbon monoxide over a 
 little pulverized cupric oxide, heated in a boat in a hard glass 
 tube (?). What gas is formed? What remains in the boat? 
 What is here the reducing agent? 
 
 112. Molecular Weight of Carbon Monoxide [Quant.]. 
 Obtain [Storeroom] a round-bottomed, 250 c.c. flask. Fit it 
 with a rubber stopper through which passes a short, straight 
 tube. Attach to the latter a short piece of rubber tubing 
 
92 CAKBON [ 113 
 
 closed with a strong pinch clamp. Make a mark on the neck 
 at the bottom of the stopper, so as to be able to measure the 
 exact content of the flask up to the stopper. Place 30 c.c. of 
 water in the flask, insert the stopper, remove the clamp, and 
 boil the water with a small flame for about five minutes, so as 
 to drive out all the air. Close the rubber tube with the clamp 
 and remove the flame quickly, wipe the flask and allow it to 
 cool. When it has assumed the temperature of the air, weigh 
 the whole carefully, suspending the flask on the balance by a 
 thread tied round the neck. Connect with the apparatus 
 delivering pure carbon monoxide, and open the clamp a very 
 little so as to admit a slow stream of the gas. When the flask 
 is full, close the clamp, disconnect from the generating appara- 
 tus, open the clamp for an instant to restore the pressure to 
 that of the atmosphere, and weigh again. The gain in weight 
 represents the weight of the carbon monoxide. Read the 
 barometer and thermometer. Subtract from the barometric 
 reading the aqueous tension at the observed temperature. 
 Ascertain the volume of the flask by filling with water to the 
 mark and weighing again. 
 
 Calculate the weight of the G.M.V. of the gas. 
 
 To what class of gases would this method of determining the 
 density and molecular weight be applicable? Why could not 
 this method be used for carbon dioxide? 
 
 113. Methane (Marsh Gas). 
 
 a. Pulverize about 5 g. of fused sodium acetate and about 
 20 g. of soda-lime (a mixture of quicklime and sodium hy- 
 droxide). Mix the ingredients intimately in the mortar, and 
 place in a large test-tube fitted with a one-hole cork and deliv- 
 ery tube (Fig. 6, p. 15). Rap the test-tube gently so as to form 
 a space for the passage of the gas. Then clamp the tube close 
 to the cork in a slightly inclined (5) position, so that it slopes 
 toward the mouth (why?). Slip over the tube a cylinder of 
 wire gauze, and, beginning at the rear or sealed end, heat the 
 contents gently until the evolution of gas is uniform. When 
 all the air is displaced, collect one bottle of gas over water. 
 
 Attach a small nozzle to the end of the delivery tube and 
 ignite the gas. Note the color and degree of luminosity of the 
 flame (?). Hold over the jet a dry inverted beaker and observe 
 one product of the combustion (?). 
 
 b. Hold the bottle containing the gas mouth downward 
 and apply a light (?). Quickly pour into the bottle some lime- 
 water, and shake (?). What volume of oxygen is required for 
 the complete combustion of one volume of methane, and what 
 
115] CAKBON 93 
 
 ratio will the resulting volume of gases bear to the original 
 volume of methane, at and at 100, respectively? 
 
 114. Ethylene [Hood]. 
 
 a. Fit a 250 c.c. flask with a doubly-bored cork, through 
 which pass a dropping-funnel and L-tube. Attach a gas- 
 washing bottle containing a little water, and connect an L- 
 shaped delivery tube. Test the apparatus to see that it is 
 air-tight. Place in the flask about 20 c.c. of phosphoric acid, 
 and clamp it to a ring-stand over a sand bath. Introduce into 
 the bulb of the dropping-funnel (or substitute, 36 b) some 
 alcohol. Finally, heat the phosphoric acid, and when it has had 
 time to reach 220, admit the alcohol drop by drop beneath the 
 surface of the phosphoric acid in the flask. 
 
 6. Fill a bottle with the gas over water and apply a light (?). 
 Quickly pour some lime-water into the bottle, and shake (?). 
 Attach a nozzle to the exit tube of the* washing-bottle and 
 raise the other tube clear of the water (why?), ignite the gas, 
 and observe the luminosity of the flame. What commercial 
 use has ethylene [R 513] ? Hold a cold, dry beaker over the 
 flame (?). What are the products of complete combustion 
 of all hydrocarbons? 
 
 c. Detach the nozzle, reattach the delivery tube to the exit 
 tube, and lower the other tube of the washing-bottle once more. 
 Place 1 c.c. of bromine [CAUTION!] in a test-tube, cover it with 
 5 c.c. of water, put the test-tube into a beaker filled with cold 
 water, and allow the gas to bubble into the bromine. When 
 the color of the bromine has disappeared, examine the contents 
 of the test-tube (?). Is the product colored? Is it soluble in 
 water? What do you infer as to its specific gravity? Detach 
 the dropping-funnel, wash it out, and place in it the contents 
 of the test-tube. Draw off the lower, oily layer into a dry 
 test-tube and note its odor (?). Try its solubility in alcohol (?). 
 
 Why is the accepted formula for ethylene preferred to the 
 simplest? What volume of oxygen would be required to burn 
 one volume of the gas completely? What would be the relative 
 volumes of the products at and 100, respectively? What 
 volumes of bromine vapor and ethylene would be required for 
 complete combination, and what would be the relative volume 
 of the product (in the state of vapor) ? 
 
 115. Acetylene. Fill a test-tube with water and invert it 
 in water in an evaporating-dish. Introduce a small piece ol 
 calcium carbide under the mouth of the tube (?). Test the 
 water with litmus (?). Bring a light to the mouth of the 
 tube (?). Note the luminosity of the flame, and compare with 
 flames of methane and ethylene (?). 
 
94 CARBON [ 116 
 
 Can you state a general method for making hydrides of non- 
 metals (such as C 2 H 2 , PH 3 , NHJ? Can SH 2 [R .645] and C1H 
 [R 715, 718] be formed similarly? 
 
 116. Illuminating -Gas. 
 
 a. Heat some sawdust in a dry test-tube (?). Note the 
 odor (?), reaction towards moist litmus paper (?), and combus- 
 tibility (?) of the vapors. What is the residue? 
 
 b. Repeat a, using bituminous coal (?). 
 
 c. Pulverize a few particles of potassium sulphate. Take a 
 small Bunsen flame and reduce the supply of air until a small 
 luminous region appears. Heat the platinum wire, touch the 
 salt with it, and hold the adhering powder steadily in the lumi- 
 nous part. After a minute or two, withdraw the bead, place it 
 upon a clean silver coin and moisten with a drop of water (?) . 
 Explain the result. What are the reducing agents used here? 
 
 117. Acids. 
 
 a. To 1 g. of sodium acetate add some dilute sulphuric acid, 
 and warm. Note the odor (?). How could you use this action 
 to make acetic acid? How is it manufactured [R 498] ? 
 
 b. Take 5 c.c. of acetic acid. Test with litmus paper (?). 
 Recall its interaction with zinc or iron (16 e) (?). To the acid 
 add 3-4 g. of litharge and boil gently for a few minutes (?). Fil- 
 ter, if necessary, while hot, and set the clear solution aside to 
 crystallize (?). What is the common name of the product [R]? 
 
 118. Alcohol. Dissolve 20 g. of glucose syrup in 150 c.c. of 
 water and add yeast. Fill a flask to the base of the neck with 
 the mixture, tie a piece of filter paper over the mouth with 
 thread [Side-shelf], and set the whole aside in a warm place for 
 3-4 days. Then warm the solution and test the gas which is 
 given off for oarbon dioxide (?). 
 
 Filter the liquid and place it in a larger flask, fitted with a 
 cork and L-tube, and connect (Fig. 14, p. 45) with a con- 
 denser [Storeroom]. Distil off about 50 c.c. Place the distil- 
 late in a distilling-flask [Storeroom] fitted with a thermometer, 
 boil with a small flame, and catch the part which passes over 
 between 75 and 93. 
 
 Note the odor of the distillate (?). Test it with litmus 
 paper (?). Use one drop to ascertain whether it burns. To 
 the rest add a crystal of iodine and enough sodium hydroxide 
 solution to dissolve it. Shake vigorously, and do not add more 
 alkali than is absolutely necessary. W^rm the solution and 
 then cool it (?). This is the iodoform test for alcohol. 
 
 119. Esters. 
 
 a. To 1 g. of sodium acetate add 1-2 c.c. of alcohol and 1 c.c. 
 of concentrated sulphuric acid. Warm, if necessary, agitate 
 
119] CARBON 95 
 
 for a minute or two, and note the odor (?). This is a test for 
 acetic acid or an acetate. 
 
 b. Place in a porcelain dish a piece of fat the size of a pea, 
 and add 4 c.c. of alcohol and five drops of a 50 per cent solu- 
 tion of sodium hydroxide (made with 1 g. of the solid and 1 c.c. 
 of water). Stir constantly and boil very gently until the odor 
 of alcohol is no longer perceptible, then stop. The alcohol is 
 used as a common solvent for the fat and the alkali. What is 
 the residue [R 505] ? 
 
 Dissolve the soap in hot water, cool, and to half of the solu- 
 tion add dilute hydrochloric acid, and shake vigorously (?). 
 Withdraw the floating coagulum by means of a glass rod, sus- 
 pend it in water in a test-tube, add a few drops of sodium 
 hydroxide, and heat until solution takes place. What do you 
 conclude from its solubility in the alkali? 
 
 To the other half of the soap solution add calcium chloride 
 solution (?). Explain the action of hard water [R 594] on soap 
 solution. 
 
CHAPTER XVI. 
 
 THE ACTIVITY OP ACIDS MEASURED CHEMICALLY. 
 
 120. Estimation of the Relative Activity of Acids. Charac- 
 terize briefly some of the methods available for measuring the 
 relative activity of acids [R 679]. 
 
 a. When methyl acetate is mixed with water, it undergoes 
 hydrolysis very slowly, acetic acid and methyl alcohol being 
 formed : CH 3 .C 2 H 3 2 + H 2 + CH 3 OH + HC 2 H 3 2 . Add about 
 1 c.c. of methyl acetate to 10 c.c. of distilled water in a test- 
 tube, test with litmus paper (?), and cork up and label the 
 mixture. After several days, test once more with litmus (?). 
 
 This action of water is found to be greatly hastened by the 
 addition of free acids, although the acids remain themselves 
 unchanged by the process. Equivalent quantities (?) of dif- 
 ferent acids show very different accelerating powers toward 
 this reaction. The order in which they are placed by measure- 
 ment of this particular form of activity, however, is the same 
 as that into which they fall when compared by any of the other 
 methods. The extent to which the change has taken place can 
 be measured at any moment by titration with alkali. The 
 quantity of acid which was present at starting being known, 
 the quantity found is the same acid plus the acetic acid set 
 free by the progress of hydrolysis. Subtraction gives the quan- 
 tity of the latter, and this quantity is a measure of the activity 
 of the accelerating acid. The activities of hydrochloric and 
 sulphuric acids are compared in ?> by this method. 
 
 6 (Two students working together). Procure two 20 c.c. 
 stoppered, graduated flasks,* and a 10 c.c. and a 1 c.c. pipette 
 [Storeroom], Mark the flasks with a file so as to be able to dis- 
 
 tinguish them, and into one measure exactly 10 c.c. of normal 
 (?) hydrochloric acid, and into the other 10 c.c. of normal (?) 
 sulphuric acid. Put exactly 1 c.c. of methyl acetate into each, 
 and fill both flasks with distilled water up to the 20 c.c. mark 
 at once (why at once?). Stopper the flasks tightly, mix the 
 contents, and suspend both so that their necks are just above 
 the water in a large bath (use the pneumatic trough) heated to 
 
 * Instead of graduated flasks, common flasks with cork stoppers 
 may be used. Measure into each flask 10 c.c. of acid and then, with 
 the same pipette, 10 c,c. of water. Add the 1 c.c. of methyl acetate, 
 shake, and proceed as directed. 
 
 96 
 
120] ACTIVITY OF ACIDS 97 
 
 about 45. If the bath is fairly large, further external heating 
 will not be necessary. Otherwise, maintain the temperature by 
 means of a small flame. In accurate work the temperature 
 must be kept constant within 0.1 during the experiment by 
 means of a thermostat. Allow the flasks to remain in this 
 position for 30 minutes (<). 
 
 While this is going on make some normal (or approximately 
 normal = 4 per cent) sodium hydroxide and fill a burette 
 with it. Take fresh portions of 10 c.c. of each of the acids and 
 titrate them with the alkali, using two drops of phenolphthalein 
 as an indicator. Record the results. These numbers measure 
 the amount of mineral acid, at starting, in each flask. 
 
 When the above time has elapsed, remove both flasks from 
 the bath, transfer the contents of each to a separate beaker, 
 rinsing out the flasks with distilled water. Add two drops of 
 phenolphthalein to each portion and titrate with the solution 
 of sodium hydroxide used before. The difference (d) between 
 the volumes of alkali required for the neutralization of 10 c.c. 
 of each acid with and without methyl acetate represents the 
 amount of sodium hydroxide required to neutralize the acetic 
 acid liberated in the hydrolysis. The two values of d obtained 
 are functions of the activity of the acids. Which acid is more 
 active? What, according to the theory of ionization, is really 
 measured in these experiments? What does the result of a 
 show water to be? 
 
 Calculate 8, the speed with constant, unit concentration, by 
 the formula for a unimolecular action [R 252] . The initial con- 
 centration (c t ) of 1:20 is equal to 0.67 moles per liter.^ The 
 proportion transformed (cc), which must be expressed in the 
 same units, is d/20 moles per liter. Then 8 = 1/t log e , 
 /c 1 /(c 1 )}. The ratio of the values of S for the two acids 
 defines their relative activities. The time (), being the same 
 for both, may be neglected, and common logarithms may be 
 used, because the factor for converting them to log e (2.3025) 
 cancels in the ratio. 
 
CHAPTER XVII. 
 
 SILICON AND BORON. 
 
 121. Silica. Mix 1 g. of finely powdered silica with 4-5 g. 
 of anhydrous sodium carbonate. Make a small watch-spring 
 spiral on the end of the platinum wire, and, by alternately heat- 
 ing in the Bunsen flame or blast-lamp, and dipping in the mix- 
 ture, obtain a large bead and heat it strongly till all action (?) 
 seems to have ceased. Place the bead in a test-tube and make 
 others by the same process. Dissolve the beads in a small 
 amount of water. Add hydrochloric acid a drop at a time 
 until the solution is strongly acid (?). Evaporate the solution 
 to dryness on the sand bath (?). Treat the residue with warm 
 water, wash the whole contents of the dish into a test-tube, and 
 examine (?). 
 
 122. Analysis of a Silicate. Mix dry potassium carbonate 
 with anhydrous sodium carbonate in equal proportions in a 
 mortar. Coil the platinum wire to watch-spring form. Mix a 
 little powdered talc (is this soluble in water? What is its com- 
 mon name?) with 6-7 times as much of the "fusion mixture," 
 and hold some* of the result on the platinum wire in the flame of 
 the blast-lamp till it is completely melted and all action (?) 
 has ceased. Repeat till several beads are obtained. Treat the 
 beads with boiling water in a test-tube until they are com- 
 pletely disintegrated. Filter through a small filter paper and 
 wash the precipitate with water. Preserve this filter paper 
 and precipitate for use later. Acidify the filtrate with concen- 
 trated hydrochloric acid and proceed as in 121. 
 
 Make a hole in the paper and wash the precipitate obtained 
 above into a test-tube. Add dilute hydrochloric acid, and 
 warm (?). Filter, if necessary, and add ammonium hydroxide 
 to alkaline reaction (?). The precipitate is aluminium hydrox- 
 ide. Boil and filter. To the filtrate add a few drops of ammo- 
 nium hydroxide, some ammonium chloride solution, and some 
 sodium phosphate solution, and shake (?). Compare 106 c. 
 
 123. Silicon Tetrafluoride. Mix intimately 1 g. of pulver- 
 ized calcium fluoride with an equal weight of sand, place in a 
 test-tube, and moisten the mixture with concentrated sul- 
 phuric acid. Apply a gentle heat (?). Hold a glass rod, with 
 a drop of water at its lower end, in the gas and examine the 
 rod (?). 
 
125] SILICON AND BORON 99 
 
 124. Boric Acid. 
 
 a. Pulverize some borax and make a strong solution in 
 boiling water. Add concentrated hydrochloric acid until the 
 solution is strongly acid, and set aside to cool (?). Filter off 
 the crystals, wash with a few drops of cold water [Note 38, 
 p. 85], and dry. Dissolve the crystals in the smallest possible 
 amount of boiling water, and set the solution aside (?). Filter, 
 and wash the crystals as before. 
 
 b. Dissolve a part of the crystals in hot distilled water. 
 Test this solution, and a sample of the distilled water, simulta- 
 neously with litmus paper (?). Dip a strip of turmeric paper 
 in the same solution, wrap it round the upper part of the test- 
 tube, and boil the solution until the paper is dry (?). Touch 
 the paper with a glass rod dipped in sodium hydroxide solu- 
 tion (?). This is a test for boric acid. 
 
 Treat the rest of the crystals with cold sodium hydroxide 
 solution (?). 
 
 c. Place on separate parts of a watch-glass a drop of concen- 
 trated sulphuric acid, a drop of glycerine, and a very little 
 pulverized borax. Rub the end of a platinum wire in each of 
 these. Bring the end of the wire slowly up to the outer edge 
 near the bottom of a small Bunsen flame. How is the flame 
 colored? This is a test for a borate. 
 
 An alternative method: Dissolve a small crystal of borax 
 in 1-2 c.c. of water in a test-tube. Add a drop or two of con- 
 centrated sulphuric acid and then 2-3 c.c. of alcohol. Heat 
 the mixture and set fire to the vapor of the alcohol. 
 
 125. Borates. 
 
 a. Dissolve 1 g. of borax in distilled water. Test both this 
 solution, and the distilled water, with litmus paper (?). 
 
 Put two drops of the solution into a test-tube and dilute 
 with water till the tube is two-thirds full. To the remainder 
 add silver nitrate solution (?). Add silver nitrate solution to 
 the very dilute solution also (?). The difference is more marked 
 if the dilute solution is first warmed. For comparison, add 
 silver nitrate solution to an exactly equally diluted sodium 
 hydroxide solution (?). What conclusion dp you draw in 
 regard to the action of water on borax? Write the equation, 
 and explain. 
 
 6. Heat a straight platinum wire in the Bunsen flame, and, 
 while yet glowing, dip it into a small quantity of borax. Re- 
 turn the wire to the flame and observe the changes in the sub- 
 stance (?) until it forms a bead. Try borax beads, as directed 
 in 3 d, e, and /, with cupric oxide (?), manganese dioxide (?), 
 and ferric oxide (?), separately. 
 
CHAPTER XVIII. 
 
 METALLIC ELEMENTS OP THE ALKALIES. 
 
 126. Potassium Hydroxide. 
 
 a. Dissolve about 30 g. of potassium carbonate [Notes 10, 
 11, 12, p. 2] in 200-300 c.c. of water in a large beaker, and 
 heat on a wire gauze to boiling. Slake 15-^0 g. of quicklime 
 in a beaker (?), using heat if necessary to start the action, and 
 make the product into a very thin paste with water. Add this 
 gradually, and with constant stirring, to the boiling solution (?). 
 Continue boiling for a few minutes (why?). Let the mixture 
 settle, and, when it is cold, decant the clear liquid (or filter 
 rapidly). Use the solution in fy c, and d. 
 
 Is calcium hydroxide appreciably soluble (Appendix IV)? 
 Is calcium carbonate more or less soluble than is the hydrox- 
 ide? Formulate the action and explain why it went to comple- 
 tion. What kind of hydroxides alone can be made by this 
 method [R 5531 ? Which hydroxides are of this kind (Appendix 
 IV)? 
 
 6. Find the strength of this solution by titration (alkali- 
 metry). To do this, place a carefully measured volume (about 
 10 c.c.) of the clear solution in a small flask. Dilute with 
 about four times its volume of water, as the concentrated 
 solution is apt to decompose the indicator. Fill a burette with 
 "normal" [R 148] hydrochloric acid* [Side-shelf], Add some 
 phenolphthalein solution to the alkali and run in the acid cau- 
 tiously until the red color just disappears. Note the volume 
 of acid used. Calculate the weight of potassium hydroxide 
 per liter, which your measurement shows to be contained in 
 the alkaline solution. Express this also in terms of a normal 
 solution containing 56 g. of KOH per liter (for example, 28 g. 
 per liter would be 0.5 normal). 
 
 c. (Two students working together). If the alkaline solu- 
 tion is above normal, calculate the volumes of water and of the 
 solution required to make 100 c.c. of normal alkali. Place the 
 necessary amount of water (by weighing) in a 100 c.c. graduated 
 
 * Or normal oxalic acid may be prepared (two students working 
 together), and used here. Calculate the weight of pure [Instructor] 
 oxalic acid [R 499] required to make 500 c.c. of normal oxalic acid 
 solution [R 148]. Weigh out [quant.] this amount on glazed paper, and 
 transfer it to a 500 c.c. graduated flask [Storeroom]. Dissolve the acid 
 in distilled water and then fill up the flask to the mark. 
 
 100 
 
127] ELEMENTS OF THE, ALKALIES 
 
 101 
 
 flask [Storeroom] and fill to the mirk ton.th ^the, 
 
 the solution is less than normal, calculate and use the "amounts 
 
 required for 200 c.c. of semi-normal alkali.) 
 
 Measure (burette) 5 c.c. of acetic acid into a flask and dilute 
 with about 20 c.c. of water (or take 25 c.c. of commercial vine- 
 gar), add phenolphthalei'n, titrate with the normal (or semi- 
 normal) alkali, and calculate the percentage of acetic acid 
 present. 
 
 d. Place very small quantities of the following solutions in 
 separate test-tubes, dilute with water, and add excess of the 
 solution of potassium hydroxide from a to each; ferric chlor- 
 ide (?) ; cupric sulphate (?) ; mercuric chloride (?). Describe the 
 color and structure of the precipitates [Note 39, below]. Boil 
 the contents of each test-tube (?). Do the precipitates dissolve 
 or change in any way? 
 
 What kind of hydroxides can be made by this method? Do 
 any metals fail entirely to form hydroxides [R 541] ? 
 
 e. Pulverize a small piece of potassium hydroxide and leave 
 it exposed to the air on a watch-glass for 24 hours, or more (?). 
 To a part of the product add dilute hydrochloric acid (?). 
 Simultaneously, treat a small piece of sodium hydroxide in 
 exactly the same way (?). Compare the results and explain. 
 
 Note 39. The structure of a precipitate may be described as 
 gelatinous, flocculent, curdy, pulverulent, granular, or crystalline. 
 What circumstances will determine the structure of a precipitate? 
 Hereafter, describe every precipitate by terms like these. 
 
 127. Potassium Nitrate. 
 
 a. Dissolve 25 g. of sodium nitrate and 22 g. of potassium 
 chloride in 50 c.c. of water and evaporate to half the volume 
 on the sand bath. Fit a filter paper into a small funnel [Note 
 24, p. 9]. As rapidly as possible decant the hot, clear liquid 
 from the crystals and set it aside. Throw the crystals which 
 appeared during boiling at once on to the filter, and rapidly press 
 out the rest of the mother-liquor with a spatula. Examine the 
 form of the crystals and ascertain what they are (?). (If they 
 are too small, recrystallize a part slowly from water in a beaker 
 in order to learn their form.) When the decanted liquid is cold, 
 press the product on a filter likewise. Examine this set of crys- 
 tals as before (?). Compare both with the original substances. 
 
 To understand the process, note the solubilities (grams dis- 
 solved bv 100 g. water) of the substances concerned (Appendix 
 V): 
 
 10 100 10 100 
 
 Potassium nitrate 246 Potassium chloride 
 
 Sodium chloride Sodium nitrate 
 
102 ELBMEFTS OF THE ALKALIES [ 128 
 
 Which cf these substances will first be deposited from the 
 boiling liquid? Ascertain by calculation how much of it 
 (roughly) will be deposited at 100, how much more will come 
 out when the liquid cools, and how much will remain in the 
 mother-liquor. What other substance will be present in large 
 quantity in the hot mother-liquor, and how much of it must 
 there be? How much of this product will be deposited when 
 the liquid cools, and how much will be lost by remaining dis- 
 solved? What per cent of the possible yield of potassium 
 nitrate may we expect to get? Dry your product, weigh it, 
 and calculate what per cent was secured. 
 
 Explain why purer potassium nitrate can be obtained by 
 crystallizing the product once more from water. 
 
 6. Mix intimately in the mortar 5 g. of finely pulverized 
 potassium nitrate with 2 g. of charcoal, and drop the mixture in 
 small portions into a red-hot crucible (?). What gases are 
 evolved? What is the residue (test with an acid)? 
 
 c. Pulverize 5 g. of potassium nitrate and mix on paper 
 [CAUTION] intimately with 2 g. of flowers of sulphur. Throw 
 the mixture [Hood] in small portions into a red-hot crucible (?). 
 What gases are evolved? Dissolve the residue and add barium 
 chloride solution (?). Explain the explosive power of gun- 
 powder (?). 
 
 128. Potassium Cyanide [POISON]. How is this salt made 
 [R 558] ? Place 2 c.c. of potassium cyanide solution in an 
 evaporating-dish, heat it, and add yellow ammonium sulphide 
 solution until the color (?) no longer disappears. Evaporate to 
 complete dryness [Hood]. Dissolve a part of the residue with 
 water, and add ferric chloride solution (?). A black precipi- 
 tate (?) indicates that the heating was not sufficient. If this 
 appears, heat the residue once more and try the action of ferric 
 chloride again. In the first part of this experiment, what 
 property of the cyanides [R 507J f and which kind of ionic chem- 
 ical change, are illustrated? Is this change instantaneous, as 
 is the union or disunion of ions? 
 
 129. Reactions of Potassium Salts. 
 
 a. Heat a little solid potassium nitrate on a clean platinum 
 wire. Notice the color of the flame and examine with the 
 spectroscope. Make a diagram showing the position of the 
 lines with reference to the D line (in the yellow), which, on 
 account of the sodium present, is shown by all flames in the 
 laboratory. 
 
 b. Saturate 10 c.c. of warm water (40) with potassium 
 nitrate. Add this solution to an equal volume of tartaric acid 
 solution in a beaker (?). Agitate and cool in a stream of 
 
131] ELEMENTS OF THE ALKALIES 103 
 
 water (?). Note, also, the effect of rubbing the inside of the 
 vessel with a glass rod. Describe the product [Note 39, p. 101]. 
 Filter, and wash the precipitate with a little alcohol. 
 
 c. Dissolve a few particles of the precipitate from b in a 
 drop or two of warm water and test with litmus (?). Remem- 
 bering that tartaric acid is a dibasic acid (H 2 C 4 H 4 6 ), what sort 
 of salt must the product be? 
 
 To a few particles of the precipitate add dilute hydrochloric 
 acid (?). 
 
 To a part of the precipitate add drop by drop potassium 
 hydroxide solution (shake between drops) (?). To the result- 
 ing solution add concentrated hydrochloric acid a drop at a 
 time. Stir vigorously with a glass rod, and cool in running 
 water between drops (?). Finally, try the effect of excess of 
 the acid (?). 
 
 Heat the rest of the precipitate strongly in a porcelain 
 crucible (?). Extract with hot water, filter, and add any acid 
 to the filtrate (?). The ignition of all potassium or sodium 
 salts of organic acids gives the same result. 
 
 d. Take 2-3 c.c. potassium chloride solution, and, in another 
 test-tube, dilute a few drops of it with 10 c.c. of water. Then 
 add some picric acid solution to each portion [Note 39, p. 101] (?). 
 Explain the difference in behavior. 
 
 What substance is shown to be present in a solution when we 
 get the tests in b and d? 
 
 130. Ammonium Salts. 
 
 a. What is a common effect of heating ammonium salts 
 (92 e)t Heat 1 g. of ammonium phosphate in a hard glass 
 test-tube (?). Dissolve the residue in water and test with 
 litmus paper (?). Is simple heating a test for an ammonium 
 salt? Do all salts of ammonium give ammonia when heated 
 (90 a; 98 a)? Observe the odors of all the salts of ammonium 
 (solids and solutions) upon the side-shelf (?), and explain. 
 
 b. Do ammonium salts impart color to the flame? 
 
 c. Take a solution of any ammonium salt and divide it into 
 three portions. To one add excess of tartaric acid solution, and 
 shake (?). To the second add picric acid solution (?). Com- 
 pare these results with 129 6 and d. To the third add sodium 
 hydroxide solution and note the odor (?). How should you 
 distinguish ammonium-ion from any other ionic substance? 
 
 d. Test ammonium chloride solution with litmus (?). Do 
 ou infer that ammonium hydroxide is a very inactive base 
 
 535, 565] ? 
 
 131. Sodium Carbonate by Solvay Process. Take 75 c.c. of 
 ammonium hydroxide, dilute with 25 c.c. of water, and dissolve 
 
104 ELEMENTS OF THE ALKALIES [ 132 
 
 in it 25 g. of powdered ammonium carbonate by shaking 
 (ammonium hydroxide solution is used instead of water in 
 order to secure ultimately a higher concentration of ammo- 
 nium-ion than could be obtained with the carbonate alone, 
 the solubility of this salt being too small). Then saturate the 
 solution completely with sodium chloride by prolonged agita- 
 tion with finely powdered salt in a corked bottle. If crude salt 
 is employed, it should be washed with water before use. 
 Decant the clear liquid into another bottle, fitted with cork and 
 two tubes, one of which reaches to the bottom. If, because of 
 delay, a dense precipitate has appeared, proceed without decant- 
 ing. Through the latter tube, pass in carbon dioxide from 
 the laboratory supply, or from a Kipp's apparatus, until the 
 solution is saturated. This operation may occupy an hour or 
 more. During the absorption of the carbon dioxide, the exit 
 tube should be closed to prevent waste of the gas. Close the 
 tubes with caps of rubber tubing plugged with glass rods and 
 set aside over night (?). Filter off the deposit, and dry by 
 pressing between filter papers. 
 
 Dissolve in water a little of the solid, which must have ceased 
 to smell of ammonia (why?), and test the reaction of the solu- 
 tion with litmus (?). If the solution is not acid, explain why it 
 is not so. 
 
 To part of the solid add any dilute mineral acid (?). 
 
 Heat the rest in a test-tube clamped so that the mouth is 
 inclined slightly downward, and ascertain what gases are 
 evolved. When gas ceases to be given off, dissolve the cold 
 residue in a very little water, test the reaction of the solution 
 with litmus paper (?), and set it aside to crystallize in an open 
 dish (?). Explain the reaction with litmus (?). Dry the 
 crystals, and ascertain the effects upon them of (a) addition of 
 an acid (?), and (b) of exposure on a watch-glass (?). 
 
 Compare the solubilities of the carbonate (Appendix IV) 
 and bicarbonate of sodium [R 544], and explain why the bicar- 
 bonate is made first and then the carbonate from it. 
 
 132. Purification of Sodium Chloride. Prepare about 150 c.c. 
 of a cold saturated solution of common salt by grinding the 
 salt for some time in a mortar with the water. If crude salt is 
 used, it must first be washed with water. Place the solution 
 in a beaker, and pass hydrogen chloride into the solution. Pre- 
 pare this gas by placing a handful of common salt in a gen- 
 erating-flask (Fig. 11, p. 29), covering it with concentrated 
 hydrochloric acid, and allowing concentrated sulphuric acid to 
 fall into it from a dropping-funnel. The hydrochloric acid 
 prevents frothing and steadies the stream of gas (why?). 
 
134] ELEMENTS OF THE ALKALIES 105 
 
 Deliver the gas into the solution through a thistle-tube with the 
 mouth downward (?). When considerable precipitation has 
 occurred, filter by putting a clean silver coin with milled edges 
 in a funnel, pouring the liquid and crystals upon it, and press- 
 ing with a spatula. 
 
 Why is the thistle-tube preferable to ordinary tubing? 
 
 133. Reactions of Sodium Salts. Take some salt of sodium 
 and try the flame test, and examine the flame with the spec- 
 troscope as in 129 a. Take a solution of some salt of sodium, 
 and add to one portion picric acid solution (?), and to the other 
 tartaric acid solution (?). Compare all the results with those 
 obtained in 129 and 130 b and c. How could you distinguish 
 sodium, potassium, and ammonium salts from one another? 
 How could you demonstrate positively the presence both of a 
 potassium and of an ammonium salt in a mixture of the two 
 (a) with and (b) without the aid of the flame test or spectroscope? 
 
 Which salts of potassium, ammonium, and sodium are least 
 soluble? 
 
 134. Ionic Equilibrium and the Ion-Product Constant. 
 
 o. Dilute a few drops of methyl orange solution [R 355] with 
 much distilled water. Add a few drops of an acid (?) and then 
 excess of a base (?). What colors does this indicator show in 
 neutral, acid, and alkaline solution, respectively? What ionic 
 substance is present in the acid, and absent from the other 
 solutions? 
 
 Take three portions of distilled water and add to each a little 
 methyl orange solution. To the first two add a little acetic 
 acid (?), to the third a drop of hydrochloric acid (?). To the 
 first add some solid sodium chloride, and shake (?). To the 
 second add some solid sodium acetate, and shake (?). What 
 ionic substance has disappeared? Explain the difference in 
 behavior [R 578]. To the third add solid sodium chloride, and 
 shake (?). Explain the absence of effect. 
 
 In what way did 63 b illustrate the same principle? 
 
 b. Take three portions of a saturated solution of potassium 
 chlorate in as many test-tubes. (If there is any deposit in the 
 bottles, this and the following solutions must be shaken before 
 use to insure saturation.) To the first add saturated sodium 
 chloride solution (?), to the second saturated potassium chlor- 
 ide solution (?), to the third saturated sodium chlorate solu- 
 tion (?). Allow them to stand fo~ a minute or two before 
 drawing any conclusion. Explain [R 581]. The experiments 
 will fail if all the solutions are not saturated. 
 
 c. Explain why, in 132, the salt is precipitated. If the crude 
 salt had contained other salts, which of them would have been 
 
106 ELEMENTS OF THE ALKALIES [ 134 
 
 affected by the introduction of hydrogen chloride, and which 
 not? Specifically, would sodium sulphate or magnesium 
 chloride have been affected, and how? If either of these would 
 be affected and had been present, under what circumstances 
 might it have entered into the precipitate (Appendix IV)? 
 Within what limits, then, does the process give a means of 
 purification? 
 
 d. To 1-2 c.c. of concentrated hydrochloric acid add con- 
 centrated sulphuric acid drop by drop (?). Explain. 
 
 Unknowns (140) may be introduced at this point. 
 
CHAPTER XIX. 
 
 METALLIC ELEMENTS OF THE ALKALINE EARTHS. 
 
 136. Calcium Oxide. Ignite [Blast-lamp] 2-3 g. of broken 
 marble for 15 minutes in an open crucible, placed upon the 
 clay triangle (?). Stir occasionally with the platinum wire. 
 When cool, add a little water (?). Test the reaction of the 
 liquid towards litmus (?). Has water any effect upon marble? 
 Now, add some dilute hydrochloric acid (?). Compare this 
 with the action of the acid upon marble (109 6). 
 
 How may the decomposition of the marble be arrested (?) 
 and reversed (?) without altering the temperature? What 
 conditions permitted complete decomposition to take place 
 here? 
 
 136. Calcium Hydroxide. Slake a piece of calcium oxide 
 and shake the product with half a liter of distilled water; let the 
 solution settle, or filter rapidly, and use the clear liquid. 
 
 a. Blow air from the lungs by means of a tube through a 
 part of the lime-water (?). How could you determine the 
 proportion of carbon dioxide in a sample of air? 
 
 6. Dilute the remainder of the lime-water with an equal 
 volume of distilled water, and pass a stream of carbon dioxide 
 from the laboratory supply, or from a Kipp's apparatus, per- 
 sistently through the solution (?). Boil a part of the resulting 
 clear liquid (?). Explain. Describe all precipitates [Note 39, 
 p. 101], both here and in all succeeding experiments. Why 
 does the second precipitate appear to differ in quantity from 
 the first? What is "temporary hardness" in water [R 594] ? 
 
 137. Reactions of Calcium Salts. Diluted calcium chloride 
 solution [Note 40, below] may be used for b, c, d, and /. 
 
 a. Try the flame test and examine with the spectroscope 
 (see that the platinum wire is clean. Note 41, below). Make 
 a sketch of the spectrum showing the positions of the lines 
 with reference to the sodium and potassium lines. 
 
 b. To a part of the solution containing calcium-ion add 
 ammonium oxalate solution in excess (?). Filter off the pre- 
 cipitate and divide it into two parts. Place each part in a test- 
 tube, and treat one with dilute hydrochloric acid (?) and the 
 other with diluted acetic acid (?). Explain this difference in 
 behavior towards hydrogen-ion [R 598]. Why was ammo- 
 nium oxalate used in preference to oxalic acid? 
 
 107 
 
108 ELEMENTS OF ALKALINE EARTHS [ 137 
 
 c. To the second portion of the solution containing calcium- 
 ion add ammonium carbonate solution (?). Warm, if necessary. 
 Filter, divide the precipitate and treat parts of it with hydro- 
 chloric acid (?) and diluted acetic acid (?) respectively. Ex- 
 plain the result. Explain also the difference in behavior of the 
 oxalate and the carbonate of calcium towards acetic acid, 
 taking account both of the solubilities of these salts, and of the 
 fundamental difference in behavior between oxalic and carbonic 
 acids. 
 
 d. To the third portion of the solution add excess of dilute 
 sulphuric acid (?). Filter and divide the clear (if not, filter 
 again) nitrate into three parts (reserve one for e). To one 
 small part add an equal volume of alcohol (formation of all 
 precipitates, if long delayed, may be hastened by vigorous 
 stirring) (?). 
 
 Neutralize (test?) the second portion with ammonium hy- 
 droxide (why?), add ammonium oxalate solution (?), and 
 explain. Is the sulphate or the oxalate of calcium more soluble? 
 Confirm your inference by giving the solubilities (Appendix 
 IV). 
 
 e. What is "permanent hardness" in water [R 595] ? Boil 
 the third portion of the filtrate from d (?). Compare the 
 result with that in 136 b. Add now a little sodium carbonate 
 solution (?). Is the carbonate or the sulphate of calcium more 
 soluble? What are their actual solubilities? How may per- 
 manent hardness be removed? What is one objection to the 
 presence of hardness in a domestic water supply (119 b, last par.) ? 
 
 /. To the fourth portion of the solution containing calcium- 
 ion add some dilute hydrochloric acid, mix, and then add 
 ammonium oxalate solution (?). Explain [R 601]. If any 
 precipitate appears, add more hydrochloric acid. Now add a 
 large amount of sodium acetate solution (?). What is the 
 precipitate (calcium acetate is soluble), and why is it formed 
 [R 601, 650] ? 
 
 Note 40. In this and all following paragraphs headed "reac- 
 tions," where "diluted" solutions are spoken of, the solutions on 
 the side-shelf must be diluted with 3-4 volumes of distilled water 
 to secure good results. 
 
 Note 41. After use, the platinum wire must be cleaned. Form 
 upon it a borax bead, and cause the molten bead to traverse the 
 wire from end to end several times. Throw off the bead. Heat 
 the wire persistently in the Bunsen flame, dipping it from time to 
 time in hydrochloric acid (why?), until the wire no longer colors 
 the flame. If much corroded, the wire may be boiled in nitric 
 icid before being heated. 
 
140] ELEMENTS OF ALKALINE EARTHS 109 
 
 138. Reactions of Strontium Salts. Use diluted [Note 40] 
 strontium chloride solution for 6, c, and d. 
 
 a. Same as 137 a. 
 
 6. Add ammonium carbonate solution (?). 
 
 c. Add dilute sulphuric acid (?). 
 
 d. Add a clear solution of calcium sulphate (made by shak- 
 ing a little of the pulverized salt with distilled water and filter- 
 ing until clear) (?). The precipitate may come slowly. Explain. 
 
 139. Reactions of Barium Salts. Use diluted [Note 40] 
 barium chloride solution for 6, c, and d. 
 
 a, b, c. Same as 138 a, b, c. 
 
 d. Add a clear solution of strontium sulphate (made by 
 shaking the salt with distilled water and filtering) (?). Explain. 
 
 Compare with 138 d, and arrange the sulphates of these three 
 metals in the order of solubility. Give two methods of dis- 
 tinguishing the compounds of the elements in this family. 
 Examine the table of solubilities (Appendix IV) and suggest 
 another way of distinguishing the three members, and one way 
 of distinguishing calcium-ion and barium-ion, respectively, 
 from the other two. How could you tell a solution containing 
 the ions of a member of this family from one containing those 
 of the members of the previous family? 
 
 140. Identification of Unknown Substances. Take three 
 dry test-tubes, apply to the instructor for three "unknown" 
 substances, and ascertain by the use of any experiments you can 
 devise what each is. 
 
 a. Study: (1) Physical appearance (?). 
 
 (2) Odor (?). 
 
 (3) Solubility in water and reaction of the solution toward 
 litmus (?). Use this aqueous solution, and employ the reac- 
 tions of the metallic ions just studied, for the purpose of recog- 
 nizing the positive radical. If the substance is insoluble in 
 water, try dilute hydrochloric acid. 
 
 (4) Effect of heating in a dry test-tube (?). Observe closely 
 the behavior of the substance, and try to identify the vapors 
 or gases given off. Preserve the residue, as, after the next 
 experiment, examination of this may be necessary. 
 
 Before trying this experiment, make a list of afl the negative 
 radicals known to you, and place opposite to each the gases, if 
 any, which salts containing that radical, when heated, might 
 be expected to give off. Consider also the means of recognizing 
 these gases, if formed. 
 
 (5) Heating with a drop or two of concentrated sulphuric 
 acid (?). 
 
110 ELEMENTS OF ALKALINE EARTHS [ 140 
 
 Before trying this experiment, add to the above list the reac- 
 tions of sulphuric acid with salts containing each of the negative 
 radicals. Consider also the means of recognizing such of the 
 possible products as are volatile. 
 
 (6) Heating the residue from (4) with concentrated sul- 
 phuric acid, or such other confirmatory experiments with this 
 residue as the results of (4) and (5) suggest (?). 
 
 (7) Other experiments suggested by the results of (l)-(6). 
 
 b. Write out the experiments and reasoning carefully in 
 your note-book, and make sure that they prove the substance 
 to be the one you finally decide that it is, and exclude the pos- 
 sibility of its being any other. Report the result to the in- 
 structor. 
 
CHAPTER XX. 
 
 COPPER AND SILVER. 
 
 141. Cuprous Chloride. 
 
 a. Dissolve about 2 g. of cupric chloride in 15 c.c. of water 
 in a small flask. Add 2-3 c.c. of pure, concentrated hydro- 
 chloric acid and about 5 g. of copper nails, and boil [Hood] 
 until the green tint is no longer perceptible in the dirty yellow- 
 ish-brown color of the product. If a few drops added to a test- 
 tube full of water confer a blue tinge upon the water, the action 
 is still incomplete. What ion confers the blue tinge? What 
 change do the cupric ions undergo? Is the change an oxidation 
 or a reduction of cupric chloride? 
 
 6. To a small part of the solution, when cold, add excess of 
 sodium hydroxide solution (?). Why is so much of this re- 
 quired? Preserve the mixture in a corked test-tube for use 
 in 143. 
 
 c. Pour the rest of the solution from a into a large amount of 
 water in a beaker (?). Expose some of the product, while 
 covered with water, to the sunlight (?). 
 
 d. To the rest of the product from c add concentrated hydro- 
 chloric acid, and shake (?). What is the complex negative ion 
 here formed [R 621]? Does it give a greater or a less concen- 
 tration of cuprous-ion than does the insoluble cuprous chlor- 
 ide? Upon this basis explain the process of solution here 
 observed. Pour a little of the solution into much water (?). 
 
 e. To the rest of the solution made in d add a little concen- 
 trated nitric acid (?). What ionic substance is shown by its 
 color to be present? Pour this solution into much water (?). 
 Consider each ionic substance originally present in the solution, 
 and explain how it is affected by the nitric acid. 
 
 Define this form of oxidation, and the form of reduction in a, 
 in terms of the hypothesis of ions. 
 
 142. Cupric Hydroxide. 
 
 a. To 1-2 c.c. of diluted cupric sulphate solution add excess 
 of sodium hydroxide solution (?). Take one-third of this mix- 
 ture and boil it (?). 
 
 b. Boil 1 g. of sugar dissolved in water (or diluted glucose 
 syrup) with a few drops of dilute sulphuric acid for several 
 minutes (?). This gives glucose and levulose [R 500]. Add 
 
 111 
 
112 COPPER AND SILVER [ 143 
 
 this glucose solution to another portion of the mixture from a, 
 warm gently (?), and note all changes. 
 
 c. To the remainder of the mixture from a add ammonium 
 hydroxide, and shake (?). What complex ion possesses this blue 
 color [R 623, 625] ? Does this ion give a greater or a less con- 
 centration of cupric-ion than does the insoluble hydroxide? 
 Upon this basis explain the process of solution here observed. 
 
 143. Cuprous Oxide. Divide the mixture from 141 b, con- 
 taining a precipitate of hydrated cuprous oxide, into three 
 parts. 
 
 a. Shake one persistently with air in a bottle, admitting 
 fresh air from time to time (?). 
 
 6. Boil the second portion (?). 
 
 c. To the third portion add ammonium hydroxide (?). What 
 is the color of the mixture? What complex ion is formed 
 [R 622]? Why is cuprous oxide dissolved by ammonium 
 hydroxide solution? 
 
 Shake this solution with air. What ion is formed? Is the 
 action more or less rapid than in a, and why? 
 
 144. Cuprous Iodide. Dilute a few drops of cupric sulphate 
 solution and add potassium iodide solution (?). Filter. Wash 
 the precipitate [Note 38, p. 85] (?). Add part of the filtrate to 
 some starch emulsion (?). Shake the rest with some chloro- 
 form (?). Read the footnote to p. 11. 
 
 145. Reactions of Cupric Salts. Use diluted cupric sulphate 
 solution. What is the color of cupric sulphate itself? To 
 what is the color of the solution due? 
 
 a. Test the reaction of the solution with litmus paper (?) 
 and explain [R 536, 344]. Is copper an active metallic ele- 
 ment? 
 
 b. Lead hydrogen sulphide through another portion (?). Is 
 this action reversible, theoretically? Add dilute sulphuric 
 acid (?). Do dilute acids act upon any sulphides (74 g)1 Why 
 not upon this one [R 600] ? 
 
 c. To another portion add potassium ferrocyanide solution 
 [R 536] (?). 
 
 d. Note here the interaction (68 a) of zinc with cupric-ion (?). 
 
 e. Make a borax bead and heat with it a minute particle of 
 cupric oxide in the oxidizing (?) and in the reducing (?) flame, 
 as in 3 e. The latter requires patience. If cupric sulphate had 
 been used here, what reaction would have taken place [R 390]? 
 
 /. Prepare a match (or splinter of wood) as in 86. Place on 
 the end a moistened mixture of any copper salt with anhydrous 
 sodium carbonate and heat in the reducing region of a small 
 Bunsen flame. Break up the charred stick gently in water in 
 
149] COPPEE AND SILVER 113 
 
 a mortar, wash away the lighter particles, and examine the 
 residue (?). 
 
 146. Ammonio-cupric Compounds. To a diluted solution of 
 cupric sulphate add ammonium hydroxide (?), first a drop or 
 two (?), then in excess (?). Of what ion does the copper now 
 form a part? Do other compounds of copper yield the same 
 ion (142 c)? Which action of all those in 146 and 146 should 
 you hold to give the most delicate test for cupric-ion? 
 
 Try with the blue solution the tests in 146 6, c, and d (?) . 
 Are the concentrations of cupric-ion given by cupric sulphide (?), 
 cupric ferrocyanide (?), cupric sulphate (?), and cupric hydrox- 
 ide (?), larger or smaller than the concentration of free cupric- 
 ion given by ammonio-cupric-ion? Which of these compounds 
 would be dissolved by ammonium hydroxide solution, and 
 which not? 
 
 147. Guprocyanides. To a diluted solution of cupric sul- 
 phate add potassium cyanide solution [CARE! POISON], first 
 a drop or two (?), then in excess (?). Does the solution show 
 the color of cupric-ion? In what form of "combination is the 
 copper [R 624] ? Divide the solution into three parts and try 
 the tests for cupric-ion given in 146 b and d (?). Would cupric 
 sulphide dissolve in potassium cyanide solution? Explain your 
 answer. To the third portion add ammonium .hydroxide (?). 
 Do cupric sulphide (?) and ammonio-cupric-ion (?) give larger 
 or smaller concentrations of free copper ions than does the 
 cuprocyanide? Would potassium cuprocyanide give the bead 
 test (146 e) or the match test (146 /)? 
 
 148. Double Salts (Potassmm-Cupric Sulphate). 
 
 a. Saturate water at 70 with 5 g. of finely powdered potas- 
 sium sulphate (about 25 c.c. will be required). Calculate the 
 weight of crystallized cupric sulphate which must be taken to 
 get an equimolecular proportion, and dissolve it in its own 
 weight of hot water. Mix the two solutions, taking care not to 
 allow any undissolved fragments of either salt to get into the 
 mixture, and set the result aside to crystallize (?). Examine 
 the form of the crystals and compare with those of blue vit- 
 riol (?). Dissolve a part of the crystals in water and use por- 
 tions of the solution for b. 
 
 b. With the solution try tests in 145 b, c, and d (?). Is cupric- 
 ion present? To another portion add ammonium hydroxide (?). 
 
 How do double salts differ (a) from complex compounds like 
 potassium cuprocyanide and ammonio-cupric sulphate and (b) 
 from simple salts like cupric sulphate? 
 
 149. Equivalent of Copper [Quant.]. Take a small rod of 
 pure zinc, smooth the ends with a file, and weigh carefully. 
 
114 COPPER AND SILVER [ 150 
 
 Place in a beaker an exactly known weight of crystallized 
 cupric sulphate (about 2 g.), and dissolve in distilled water. 
 Put the zinc in this solution, and allow them to remain in con- 
 tact until the latter is completely decolorized. Remove the 
 zinc, free it carefully from the brown deposit (?), and dry and 
 weigh it. What weight of zinc has gone into solution? 
 
 To avoid weighing the precipitate of copper, which it would 
 be difficult to do exactly, calculate from the formula what 
 quantity of copper was contained in the amount of blue vitriol 
 taken (?). Calculate from your data the weight of copper dis- 
 placed by the equivalent weight of zinc found in 35 b or 36 a, or, 
 if zinc was not then employed, assume the latter to be 32.7. 
 This weight of copper will be the equivalent (that of oxygen 
 being 8). Look up the specific heat of copper (Appendix III), 
 and use this and the equivalent observed to find the atomic 
 weight (?). 
 
 What other atomic weights could be measured on this plan? 
 
 150. Reactions of Silver Salts. 
 
 a. Take 1-2 c.c. of silver nitrate solution. Test its reaction 
 towards litmus [R 536]. (?). Is silver more or less active as a 
 metallic element than copper? 
 
 Add dilute hydrochloric acid until no further precipitation 
 occurs (?)'. Filter, wash the precipitate with water, and place 
 a small part of it in the sunlight (?). 
 
 What effect does the skin have on silver nitrate? 
 
 b. To a small part of the precipitate add ammonium hydrox- 
 ide (?). What complex ion is formed ? Now add dilute nitric 
 acid in excess (?). Formulate the action of this acid. 
 
 c. To another small part add sodium thiosulphate solution (?) 
 Pass hydrogen sulphide (see /) through the solution (?). Ex- 
 plain both actions. 
 
 d. Place the rest of the precipitate in a porcelain crucible, 
 put on it a piece of granulated zinc, and fill up with dilute sul- 
 phuric acid. Stir from time to time (?). After an hour or two 
 pour off the acid, take out any unchanged zinc, wash the pre- 
 cipitate with water by decantation, add ammonium hydroxide, 
 and filter. Find out whether any silver chloride had remained 
 unchanged and so passed into the filtrate (150 6) (?). When 
 the filter paper is dry, place the dark powder in a hollow on a 
 stick of charcoal and melt it with the flame of the blast-lamp 
 directed downward upon it (?). 
 
 e. Dilute a few drops of silver nitrate solution and divide 
 into two parts. To one add potassium bromide solution (?), 
 and to the other potassium iodide solution (?). Add to each 
 some ammonium hydroxide (?). Compare the rates of action 
 
150] COPPER AND SILVER 115 
 
 with that on silver chloride (?). Arrange the three salts and 
 ammonio-argentic-ion in the order of decreasing ability to give 
 argentic-ion (?). Relate this order to that of solubility (Appen- 
 dix IV). 
 
 /. Dilute 2-3 drops of silver nitrate solution and lead in 
 hydrogen sulphide (?). Is this action reversible, theoretically? 
 Divide the product into two parts, and to one add dilute nitric 
 acid (?). 
 
 To the other part add potassium cyanide solution [CAUTION! 
 POISON] (?). What is formed? Could silver chloride be pre- 
 cipitated from this solution (64 6)? Explain. 
 
 g. To 1 c.c. of silver nitrate solution add a few drops of 
 potassium dichromate solution (?). Test the solutions before 
 (?) and after mixing (?) with Congo red paper [R 355]. (If 
 the color of the dichromate obscures that of the Congo red, 
 wash the test-paper with distilled water.) What product does 
 this show? Make the equation accordingly. 
 
CHAPTER XXI. 
 
 MAGNESIUM, ZINC, CADMIUM, MERCURY. 
 
 161. Magnesium. Mix thoroughly in a mortar equal bulks 
 of magnesium powder and powdered calcium carbonate. Put 
 the mixture in a test-tube (it should fill about half an inch of 
 the tube), fix the tube in a clamp on the stand, and heat the 
 top layer in the Bunsen flame until the reaction begins. Be 
 careful to keep the tube directed away from the face during the 
 heating. Allow the test-tube to cool, add a little water, and 
 then, slowly, an excess of concentrated hydrochloric acid (?). 
 (If the tube has been broken, place the contents with the acid 
 in a beaker.) What effect will the acid have upon any excess 
 of either of the ingredients? The acid will also dissolve the 
 oxides of magnesium and calcium formed by the action. When 
 all action has ceased, filter and wash the black residue (?) with 
 water. After drying this on the radiator or water bath, prove 
 that it is carbon. This may be done by placing some of it in a 
 dry test-tube, adding a pinch of potassium chlorate, heating in 
 the Bunsen flame, and pouring the gas when it has cooled (close 
 the tube with the thumb while waiting for this) into a test- 
 tube containing 2 c.c. of lime-water, and shaking (?). 
 
 What is the reducing agent in this action? 
 Examine the test-tube (?) and explain [R 518]. . 
 
 162. Properties of Magnesium Compounds. 
 
 a. Try whether magnesium chloride dissolves completely in 
 water (?). Test the solution with litmus (?). 
 
 Heat some of the crystals strongly in a dry test-tube (?). 
 Test the reaction towards litmus paper of the water which con- 
 denses in the tube (?), and then remove the liquid from the 
 sides of the tube with a piece of filter paper. Does the residue 
 dissolve in water? Explain. 
 
 b. To some diluted magnesium sulphate solution add ammo- 
 nium hydroxide (?). Explain the result in terms of the ion- 
 product constant [R 587] (?). Now mix with some ammonium 
 hydroxide several times its volume of ammonium chloride solu- 
 tion (what effect will this have on the ionization of ammonium 
 hydroxide? Would ammonium sulphate answer as well?), and 
 then add the mixture to a new portion of magnesium sulphate 
 solution (?). Explain in terms of the ion-product constant. 
 
 To this combination of three solutions add sodium phosphate 
 
 116 
 
153] MAGNESIUM, ZINC, MERCURY 117 
 
 solution (see 106 c) (?). Explain the purpose of each ingre- 
 dient. 
 
 c. To a fresh portion of the diluted magnesium sulphate 
 solution add ammonium carbonate solution, and warm (?). 
 Repeat, adding excess of ammonium chloride solution to the 
 magnesium sulphate solution before using the carbonate (?). 
 
 Calcium-ion (137 c), strontium-ion (138 6), and barium-ion 
 (139 6) were also precipitated by ammonium carbonate. 
 Repeat these experiments with them, adding first excess of 
 ammonium chloride solution (?). If you had a salt of magne- 
 sium mixed with a salt of one of those other metals, how should 
 you proceed so as to precipitate a compound of the alkaline 
 earth metal first and one of magnesium afterwards? 
 
 Add two drops of hydrochloric acid (why?) to about 250 c.c. 
 of the city water, evaporate to small bulk, and test it for cal- 
 cium-ion and magnesium-ion. 
 
 d. Pass hydrogen sulphide through magnesium sulphate 
 solution (?). 
 
 153. Reactions of Zinc Salts. Use diluted zinc sulphate 
 solution. 
 
 a. Test the solution with litmus paper (?). Explain. 
 
 6. To a part of it add sodium carbonate solution (?), at first 
 a little and then in excess. Note the gas evolved (?). Why is 
 the gas slow in appearing? Relate this result to that in a (?). 
 Bring the contents of the tube to the boiling-point, filter, and 
 wash the precipitate with water. To a portion of the precipi- 
 tate add an acid (?). Account for the evolution of gas during 
 the precipitation. Dry the rest of the precipitate for c. 
 
 c. Heat the dried, basic zinc carbonate from b in a porcelain 
 crucible to redness for a few minutes. Remove a small portion 
 of the product and try the action of an acid upon it (?). If it 
 effervesces, ignite for a longer time. What is the color of the 
 product when hot, and when cold? Reserve for d. 
 
 d. Moisten the residue from c with a few drops of a cobalt 
 chloride solution and heat again (?). 
 
 e. To another part of the zinc sulphate solution add a very 
 little sodium hydroxide solution, and shake (?). Filter, sus- 
 pend the precipitate in water, and divide into three parts. To 
 one add an excess of sodium hydroxide solution (?). Does this 
 show zinc hydroxide to be basic, or acidic? To the second por- 
 tion add dilute hydrochloric acid (?). What sort of hydroxide 
 does it now seem to be? Explain [R 648]. 
 
 To the third portion add ammonium hydroxide (?). What 
 complex ion is formed [R 648] ? 
 /. To a third portion of zinc sulphate solution add ammo- 
 
118 MAGNESIUM, ZINC, MERCURY [ 154 
 
 nium sulphide solution (?). Filter, and preserve the precipi- 
 tate for g. 
 
 g. Roll up a small part of the filter paper from / into a ball, 
 and coil the platinum wire tightly round it. Roast the whole 
 in the Bunsen flame (?). Moisten the ash with cobalt chloride 
 solution and heat again (?). 
 
 154. Relative Activity of Several Acids. In clean test-tubes 
 place equal volumes of (1) zinc chloride, (2) zinc sulphate, and 
 (3) zinc acetate solutions. Compare their reactions towards 
 litmus and towards Congo red paper [R 356] (?). Into each 
 pass hydrogen sulphide to saturation (test? Note 36, p. 67) (?). 
 Are actions like this reversible, theoretically? The reverse 
 action consists in the action of what acid upon what insoluble 
 salt in each case? Will it be equal with different acids [R 599]? 
 If not, the most active acid will have kept the most zinc in solu- 
 tion and the least active the least. To ascertain how much 
 zinc remains in each solution, filter the mixtures separately, and, 
 after testing with Congo red paper (?), add ammonium hydrox- 
 ide to each (?). The precipitates are zinc sulphide (why?), 
 Compare the amounts (?). Infer the relative activities of the 
 acids (?). 
 
 Confirm the conclusion, roughly, by putting some of the 
 precipitate into each of three test-tubes, treating directly with 
 dilute hydrochloric, sulphuric, and acetic acids, of equivalent 
 concentrations, and comparing the rates of action (?). 
 
 Why was ammonium sulphide used in 153 / ? 
 
 155. Ionic Equilibrium. Take a larger amount of zinc sul- 
 phate solution, and add sulphuric acid to it cautiously until a 
 sample just ceases to give any precipitate with hydrogen sul- 
 phide. Explain. Now add much powdered, anhydrous sodium 
 sulphate, stir until it has dissolved, and test a part with hydro- 
 gen sulphide again (?). What effect must the great addition of 
 sulphate-ion have upon the hydrogen-ion introduced by the 
 sulphuric acid? Why is zinc sulphide now precipitated? 
 
 156. Reactions of Cadmium Salts. Use diluted cadmium 
 sulphate solution. 
 
 a. Same as 153 e. Answer the same questions. 
 
 b. Saturate with hydrogen sulphide (?). Is the action easily 
 reversible? Add dilute hydrochloric acid (?). 
 
 By what reactions could you distinguish between salts of 
 magnesium, zinc, and cadmium? 
 
 157. Unknown Substances. Apply to the instructor for 
 three unknown substances. To identify the anions, follow 
 the outline of work in 140. To identify the cations, employ 
 
161] MAGNESIUM, ZINC, MERCURY 119 
 
 the scheme commonly used in analysis [R 660]. Present to the 
 instructor a logical and coherent, written report. 
 
 168. Mercurous Nitrate. Place about 10 g. of mercury with 
 15 c.c. of diluted (1 : 1) nitric acid in a small beaker, and let the 
 action go on for an hour. Stirring will cause crystallization. 
 Dissolve the crystals in water to which a few drops of nitric acid 
 have been added (why?). Use this solution in 160. 
 
 169. Mercuric Nitrate. Boil some mercury or mercurous 
 nitrate with excess of concentrated nitric acid. Evaporate the 
 solution on a water bath, moisten with nitric acid, and dry. 
 
 160. Reactions of Salts of Mercury. Use diluted portions of 
 mercurous nitrate solution and of a diluted solution of any 
 mercuric salt, and add to each the following reagents. Com- 
 pare the results in each case. 
 
 a. Litmus paper (?). Explain. 
 
 b. Dilute hydrochloric acid (?). Treat the precipitate, if 
 there is any, with ammonium hydroxide [R 659] (?). 
 
 c. Sodium hydroxide solution [R 656] (?). 
 
 d. Ammonium hydroxide [R 659] (?). 
 
 e. Hydrogen sulphide to saturation [R 657] (?). 
 
 /. Potassium iodide solution (shake) until there is no further 
 change (?). What complex ion is formed? Try an experiment 
 to show that mercuric sulphide gives an even smaller concen- 
 tration of mercuric-ion than does this complex ion (?). 
 
 g. Stannous chloride till there is no further change [R 655] (?) . 
 
 h. A clean copper nail (?). Determine whether any copper 
 goes into solution (?). Explain (Appendix VII). 
 
 i. Heat a little of a salt of mercury in a dry test-tube (?). 
 
 /. How could you distinguish a solution of a mercurous and 
 of a mercuric salt, respectively, from salts of silver, copper, 
 magnesium, zinc, and cadmium? 
 
 Do magnesium, zinc, cadmium, and mercury exhibit the 
 properties of typical metallic elements [R 533]? Explain. 
 
 161. Concentration Cell. Suspend a rod of tin about 60 mm. 
 long by a thread from one end, and hang it near the bottom of 
 the graduated cylinder. Pour in through the dropping-funnel, 
 which must reach the bottom of the cylinder, first highly 
 diluted, dilute hydrochloric acid (1 : 6 Aq.) and then diluted 
 (1 : 1) stannous chloride solution. Perform the operation with 
 care, in such a way that the solutions do not mix, the dilute 
 acid being finally uppermost, and that the surface at which 
 they meet is near the middle of the rod of tin. If the second 
 solution is permitted to carry air bubbles with it, mixing will 
 inevitably occur. Place the arrangement where it will not be 
 disturbed, and examine from time to time (?). Explain [R 673]. 
 
CHAPTER XXII. 
 
 ALUMINIUM, TIN, LEAD. 
 
 162. Aluminium. 
 
 a. Record here the action of aluminium on hydrochloric acid 
 (16 a) and its degree of intensity (?). Try the action of the 
 metal on dilute sulphuric (?) and dilute nitric acids (?). Com- 
 pare and explain. 
 
 b. Heat some aluminium turnings with sodium hydroxide 
 solution for some minutes (?). To ascertain whether anything 
 has gone into solution, neutralize one-half of the liquid care- 
 fully with dilute hydrochloric acid (?). Test the precipitate 
 by!63/(?). 
 
 c. To the rest of the liquid from b add calcium chloride solu- 
 tion (?). To what group of minerals do compounds allied to 
 this belong [R 685] ? 
 
 163. Reactions of Aluminium Salts. Use portions of diluted 
 aluminium sulphate solution. 
 
 a. Test the solution with litmus paper (?). Explain. 
 
 b. Add ammonium sulphide solution (?). Filter off the 
 precipitate, wash it until odorless (why?), and ascertain whether 
 it is a sulphide or not (?). Preserve a part of it for /. 
 
 c. Add sodium carbonate solution (?). Filter off the precip- 
 itate, wash it until free from sodium carbonate (why?), and 
 ascertain whether it is a carbonate or not (?). 
 
 d. Add a little sodium hydroxide solution (?). Proceed as 
 in 153 e, and answer the same questions (?). Explain the dif- 
 ference in behavior observed (?). 
 
 e. To a small portion add sufficient sodium hydroxide solu- 
 tion to redissolve the precipitate. Then add excess of ammo- 
 nium chloride solution, and boil (?). Explain. 
 
 /. Proceed as in 153 g with a small part of the filter paper 
 from b (?) (or use the method in 153 d). 
 
 164. Alum. Prepare warm, saturated solutions of hydrated 
 aluminium sulphate and ammonium sulphate in approximately 
 equimolecular proportions (calculate amounts required), mix 
 them, and set aside (?). Obtain some large crystals by hanging 
 a thread in the solution. Note the form and taste of the 
 crystals (?). 
 
 Dissolve one or two of the crystals, and ascertain by experi- 
 ments selected from 163 whether the solution contains alu- 
 
 120 
 
168] ALUMINIUM, TIN, LEAD 121 
 
 minium-ion (?). Infer whether this is a double salt, or a salt 
 giving a complex ion. 
 
 165. Mordanting. 
 
 a. To some cochineal solution add any solution containing a 
 salt of aluminium and then ammonium hydroxide (?). Filter. 
 What becomes of the coloring matter? 
 
 b. Soak a small piece of cloth in a strong solution of alum 
 (made in 164), and then transfer it to a beaker containing a 
 boiling solution of cochineal made strongly alkaline with ammo- 
 nium hydroxide (?). What purpose does the alum serve? 
 What is the object of the ammonium hydroxide? What com- 
 pounds of aluminium would act as mordants without the addi- 
 tion of a base (?), and how do they so act [R 689] ? 
 
 166. Tin. Place some tin (gran.) in a test-tube with diluted 
 nitric acid (1 acid: 10 Aq) and set aside. After 168, examine 
 portions of the solution. Determine whether it contains a 
 salt of ammonium (?), and explain. Determine whether it 
 contains stannic-ion or stannous-ion (?). How does tin behave 
 with concentrated nitric acid (97 e) ? Explain the difference. 
 
 167. Halides of Tin. 
 
 a. Stannous halide. Heat about 1 g. of tin with pure, con- 
 centrated hydrochloric acid (see 16 a) (?). Let the action go 
 on until the acid is nearly exhausted. Use the solution in 6, 
 and in 168. Proceed with later experiments until it is ready. 
 
 b. Stannic halide [Hood]. To one-half of the solution from a 
 add bromine-water until the color ceases to be destroyed, and 
 drive off the excess of bromine by warming (?). Use this 
 liquid in 168. 
 
 168. Reactions of Stannous and Stannic Salts. In a, b, and 
 c use a portion of each of the solutions from 167, after clilution, 
 with each reagent. Compare the two results in each case. 
 
 a. Saturate (test?) each portion with hydrogen sulphide (?). 
 To part of each product add dilute hydrochloric acid, to learn 
 whether the action is easily reversible (?). Filter the remainder 
 of each product, and treat the precipitates, separately, with 
 warm yellow ammonium sulphide solution [R 697] (?). To 
 the resulting liquids add dilute hydrochloric acid (?). Explain 
 why both give the same product (?). What is the gas evolved? 
 
 b. Add mercuric chloride solution, at first a little and then 
 in excess, first to the stannic (?) and then to the stannous solu- ' 
 tion (?). When the changes (two) in the latter are complete, 
 boil, let the precipitate settle, filter the clear part of the liquid, 
 and determine whether the tin in the filtrate is now stannous or 
 stannic by adding a drop or two of bromine-water (see 167 b). 
 
 c. Add to each portion a little sodium hydroxide solution (?). 
 
122 ALUMINIUM, TIN, LEAD [169 
 
 Proceed as in 158 e and answer the same questions (?). Ex- 
 plain the difference, if any, in behavior observed. What other 
 hydroxide resembles those of tin and zinc? 
 
 d. Boil a small portion of the stannic solution with tin (gran.) 
 for several minutes. Now add mercuric chloride solution (?) 
 and compare with the results in b (?). Account for the change 
 in the stannic-ion (?). What kind of chemical change was this ? 
 
 169. Apply to the instructor for two unknown substances. 
 Identify them, and report the result, as directed in 140 and 157. 
 
 170. Lead. 
 
 a. Dissolve 1 g. of lead acetate in 20 c.c. of water, place in 
 it several pieces of granulated zinc, and set aside for an hour or 
 two. After 171 devise a way of precipitating any remaining 
 lead-ion and showing the presence of zinc-ion in the solution. 
 
 b. Wash some of the lead from a with distilled water, and see 
 whether it is possible to get washings which show no reaction 
 with hydrogen sulphide (?). Account for what you observe. 
 
 171. Reactions of Lead Salts. Use diluted lead nitrate. 
 a. Test with litmus paper (?). 
 
 6. Saturate (test?) with hvdrogen sulphide (?). Is the 
 action easily reversed? 
 
 c. Add dilute hydrochloric acid (?). Filter, dilute the fil- 
 trate with 5-10 volumes of water, and saturate with hydrogen 
 sulphide (?). Explain. What other chlorides are "insoluble"? 
 
 d. Add sodium hydroxide (?), first a little, then in excess. 
 What other hydroxides resemble this one? 
 
 e. Add potassium iodide solution (?). Boil, filter, and 
 examine the filtrate (?). Infer a property of lead iodide (?). 
 
 /. Potassium dichromate solution (?). Proceed as in 150 g (?). 
 g. Dilute sulphuric acid (?). What sulphates are insoluble? 
 
 172. Lead Dioxide. 
 
 a. Warm and agitate 1 g. of minium with 5-6 c.c. of dilute 
 nitric acid until it no longer changes in color (?). Dilute with 
 water, and filter. Reserve the precipitate for b and c. Show 
 by tests selected from 171 that lead-ion is present in the fil- 
 trate. What theory of the nature of red lead is suggested by 
 this action [R 701] ? 
 
 b. Treat a part of the precipitate from a with sodium hy- 
 droxide solution (?). Record here the action of hydrochloric 
 acid upon lead dioxide (29 d) and explain it (?). 
 
 c. Allow the rest of the precipitate from a to dry, place it 
 in an evaporating-dish, and direct upon it a stream of hydrogen 
 sulphide (?). Explain. 
 
 173. Do aluminium, tin, and lead exhibit the properties of 
 typical metallic elements [R 533] ? Explain. 
 
CHAPTER XXIII. 
 
 ARSENIC, ANTIMONY, BISMUTH. 
 
 174. Arsenic. 
 
 a. Heat one particle of arsenic in a hard glass test-tube (?). 
 
 b. Roast a particle of arsenic on a crucible lid (?). Note the 
 behavior and odor (?). 
 
 c. Mix a pinch of pulverized arsenic trioxide with wood 
 charcoal (pulv.). Heat the mixture strongly in a dry test- 
 tube (?). Compare result with a (?). 
 
 d. Boil 0.5 g. of arsenic (pulv.) with excess of nitric acid (?) . 
 Compare with 105 b. Use the solution for 177 e. 
 
 175. Arsenic Trioxide and Arsenious Chloride. 
 
 a. Boil 0.5 g. of arsenic trioxide with water (?). Test the 
 solution with litmus (?). Now add sodium hydroxide solution, 
 and boil (?). To what class of oxides does the trioxide appear 
 to belong? Formulate the whole change. 
 
 b. Boil 1 g. of the trioxide with 2-3 c.c. of concentrated 
 hydrochloric acid (?). To what class of oxides does it appear 
 now to belong? Formulate this change. Set aside and ex- 
 amine later (?). What are the crystals [R 710] ? Is this action 
 reversible? Use the solution for c and 178. 
 
 Is arsenic a typical metallic element? Is it a metal at all? 
 Give reasons taken from a and b. 
 
 c. Take a small part of the solution from b (setting aside 
 the rest), dilute with water, and saturate with hydrogen sul- 
 phide (?). Divide into two parts. With one, try whether the 
 action is easily reversible (?). Filter the other and heat the 
 precipitate with yellow ammonium sulphide solution (?). What 
 other sulphide behaved in this way? Now add dilute hydro- 
 chloric acid to the liquid (?). Explain. What is the gas 
 evolved? 
 
 176. Arsenites. Use portions of diluted potassium arsenite 
 solution. 
 
 a. Add silver nitrate solution (?), then ammonium hy- 
 droxide (?). 
 
 b. Add cupric sulphate solution [R 712] (?). 
 
 c. Test the solution with litmus (?). What substances must 
 be present? Saturate with hydrogen sulphide (?). 
 
 177. Arsenic Acid and Arsenates. For a, b, c, and d, use 
 portions of diluted potassium arsenate solution. 
 
 123 
 
124 ARSENIC, ANTIMONY, BISMUTH [ 178 
 
 a. Add silver nitrate solution (?). Add now ammonium 
 hydroxide (?). 
 
 b. Add magnesia mixture (106 c) (?). Compare with 152 
 b (?). 
 
 c. Add to ammonium molybdate solution as in 106 c (?). 
 
 d. Add 2-3 drops of dilute hydrochloric acid (?) and saturate 
 with hydrogen sulphide (?). Now add excess of concentrated 
 hydrochloric acid and saturate again (?), heating to assist the 
 action. Explain. 
 
 e. Neutralize the solution from 174 d with ammonium 
 hydroxide, avoiding excess, and use tests selected from a, b, 
 and c (?). 
 
 178. Arsine [HooD. CARE! POISON]. Arrange a side- 
 neck test-tube (or small flask) with safety and straight delivery 
 tubes and nozzle to generate and burn hydrogen. Place in it 
 a piece of chemically pure zinc, and add pure hydrochloric acid 
 [Side-shelf]. When the air has been displaced (CARE. Test?), 
 light the gas and hold a crucible lid in the flame (?). If there 
 is no deposit, add a drop or so of the solution of arsenic tri- 
 chloride (176 b}, observe the appearance of the flame, and 
 obtain a deposit on the crucible lid (?). What kind of chemical 
 change takes place in the flame (see 72 6)? Heat the tube, 
 through which the gas passes to the nozzle, with a Bunsen 
 flame (? Marsh's test). When these experiments are com- 
 pleted, fill the test-tube with water to stop the action. 
 
 Apply bleaching powder solution (fresh) to the deposit (?). 
 
 Can arsine be made according to the general method con- 
 sidered in 116? What other hydrides behave like arsine when 
 heated? 
 
 179. Antimony. 
 
 a, b. Proceed as in 174 a and b, using antimony (?). 
 
 180. Antimony Trioxide. 
 
 a. To obtain the trioxide, heat 2-3 g. of pulverized antimony 
 with concentrated nitric acid in a small flask [Hood] (?). How 
 should you determine whether the product was a nitrate or 
 not? (?). What other metal behaved similarly when treated 
 with nitric acid? Compare with arsenic (174 d and 177 e) (?). 
 Dilute, filter, wash well [Note 38, p. 85], and use in 6, c, d, and e. 
 
 b. Boil a small part with water (?). Add sodium hydroxide 
 solution, and boil (?). What kind of oxide is it? 
 
 c. Boil a small part with hydrochloric acid (?), What kind 
 of oxide is it? 
 
 d. Boil a part with an equal amount of potassium-hydrogen 
 tartrate in 5-6 c.c. of water (?). Filter, and set the filtrate 
 aside (?). What are the crystals? 
 
186] ARSENIC, ANTIMONY, BISMUTH 125 
 
 e. Boil the rest persistently with 2-3 c.c. of concentrated 
 nitric acid (?). Evaporate [Hood] the clear liquid, and heat 
 the residue a little below a red heat [R 715] (?). 
 
 181. Antimony Trichloride and Trisulphide. 
 
 a. Treat about 0.5 g. of antimony trichloride with water (?), 
 and test the liquid with litmus paper (?). Add more water, 
 warm, and ascertain whether the action is reversible by adding 
 drop by drop (shake between drops) concentrated hydrochloric 
 acid (?). When the liquid has become clear add a large amount 
 of water (?). What law is illustrated? Finally, add concen- 
 trated hydrochloric acid again (?), and use the solution in 6. 
 
 How does this action differ from that of water upon phos- 
 phorus trichloride (107 6) and upon arsenic trichloride [R 710]? 
 What is the significance of this difference? 
 
 6. Dilute the liquid from o, saturate with hydrogen sulphide 
 (?), and proceed as in 175 c (?). Answer the same questions. 
 
 182. Stibine. Repeat 178, using antimony trichloride (?). 
 
 183. Bismuth. Prepare a match (or splinter of wood) as in 
 86, and proceed as in 145 /, using any bismuth salt (?). What 
 metals are obtainable in this precise way? Could zinc, mercury, 
 silver, and aluminium be so obtained? Explain. 
 
 184. Compounds of Bismuth. 
 
 a. Warm about 1 g. of bismuth with 8-10 c.c. of diluted 
 (1 : 1) nitric acid (?). Concentrate to 3 c.c. and set aside 
 [R 718] (?). Compare this result with arsenic (174 d) and 
 antimony (180 a), and interpret (?). 
 
 6. Proceed as in 181 a (first par.), using bismuth nitrate (or 
 the product from a) and nitric acid, instead of the salt and 
 acid there employed (?). 
 
 c. Take 1-2 c.c. of bismuth nitrate solution, dilute it, and 
 clear up with nitric acid if necessary. Saturate a part with 
 hydrogen sulphide (?), and proceed as in 175 c (?). Compare 
 with the sulphides of arsenic and antimony (?). 
 
 d. To the remainder of the solution of bismuth nitrate add 
 sodium hydroxide (?), at first a little, and then in excess (?). 
 Compare this hydroxide with the oxides of arsenic (175 a) (?) 
 and of antimony (180 6) (?). 
 
 Filter off the precipitate, ignite it in a porcelain crucible, 
 and note the color (?) when hot (?) and when cold (?). 
 
 185. Do arsenic, antimony, and bismuth exhibit the prop- 
 erties of typical metallic elements? If not, in what respects do 
 they fail to do so? 
 
 186. Apply to the instructor for two unknown substances. 
 Identify them, and report the result, as directed in 140 and 157. 
 Do not use Marsh's test for arsenic. 
 
CHAPTER XXIV. 
 
 CHROMIUM, MANGANESE. 
 
 187. Chromates. Melt 5 g. of potassium carbonate with 
 equal amounts of potassium hydroxide (omit this from the 
 equation) and potassium nitrate (include only oxygen, from this, 
 in the quation) at a low temperature in an iron :rucible [Store- 
 room] and stir in (use the reverse end of a file) 5 g. of powdered 
 chromitc. Heat strongly [Blast-lamp] for several minutes (?). 
 When the mass has cooled dissolve it in a little boiling water. 
 Filter, and add dilute nitric acid to the solution until it is acid (?). 
 Note the change in color (?). 
 
 188. Bichromates and Chromates. Take some potassium 
 dichromate solution and run into it potassium hydroxide solu- 
 tion from a burette till the change in color is complete. A 
 test-tube trial will show the tint to be reached. Concentrate 
 the solution and allow it to crystallize (?). What kind of salt 
 (neutral, acid, basic, double, or complex) is potassium dichro- 
 mate essentially? What are the colors of dichromate-ion and 
 of chromate-ion? 
 
 189. Chromic Anhydride. Make a cold saturated solution of 
 5 g. of sodium dichromate, add to it two volumes of concen- 
 trated sulphuric acid in a beaker, and cool (?). Filter through 
 a small plug of asbestos, and dry the precipitate by smearing 
 it on a piece of broken bisque plate [Storeroom]. 
 
 190. Chromic Oxide. Pulverize potassium dichromate (10 
 g.) thoroughly with one-fifth its weight of sulphur, and heat 
 with the blast-lamp in a porcelain crucible for fifteen minutes. 
 Grind up the resulting mass in a mortar with water, filter, 
 wash the green residue (?), and dry it on a radiator for use 
 in 191. 
 
 Make a borax bead, dissolve a particle of chromic oxide in 
 it, and note the effects of the oxidizing and reducing flames 
 upon it 0). All chromium compounds give the same result. If 
 chromic sulphate had been used, what would have been the 
 nature of the chemical action? 
 
 101. Chromic Chloride. Mix the chromic oxide prepared in 
 190, with one-third its weight of pulverized wood charcoal, 
 make into a stiff paste with some starch, and mold the mixture 
 into little pellets of the size of peas. Coyer these completely 
 with a layer of charcoal powder (why?) in a closed crucible. 
 
 126 
 
195] CHROMIUM, MANGANESE 127 
 
 dry them by heating gently with the Bunsen flame, and let them 
 cool before exposing them to the air (why?). Place them in a 
 piece of hard glass tubing. Then connect with a chlorine 
 apparatus, and, when the chlorine gas has reached the pellets 
 and completely displaced the air (why?), heat strongly with a 
 blast-lamp. Conduct any superfluous chlorine into a test- 
 tube filled with sodium hydroxide. Describe the substance 
 which is formed, and try its solubility in water and acids. 
 
 192. Chrome-Alum. Dissolve 10 g. potassium dichromate in 
 water, add the amount (calculated) of sulphuric acid necessary 
 to form potassium sulphate and chromium sulphate, warm and 
 add alcohol (7-10 c.c.), a little at a time, until the yellow color 
 has entirely given place to a pure, bright green. The action 
 takes some time to reach completion. Note the odor (?). Set 
 the greater part of the solution aside to evaporate spontane- 
 ously. Concentrate the smaller portion on the water bath until 
 crystals appear. Examine the form and color of the crystals 
 from both portions (?). What is the color of their solution in 
 water? What is the color of chromic-ion? 
 
 193. Reactions of Chromic Salts. Make a solution of chrome- 
 alum. What are the ions in the solution? 
 
 a. Boil a portion for some time [R 730]. 
 6. To another portion add sodium hydroxide solution, at 
 first a little (?), then in excess (?). Boil [R 729]. 
 
 c. Add ammonium sulphide (?). Filter off the precipitate, 
 wash it until odorless, and determine whether it is a sulphide. 
 
 d. Add excess of sodium hydroxide solution and then a large 
 volume of bromine-water, and heat (?). Try another portion, 
 using lead dioxide instead of bromine (?). Infer the nature of 
 the action from the change in color. 
 
 194. Reactions of Chromates. For a, 6, c, d, use diluted po- 
 tassium chromate solution. What are the ions in the solution? 
 
 a. Acidify a part of the solution with dilute sulphuric acid (?). 
 Concentrate and set aside to crystallize (?). 
 
 6. Recall the actions of hydrogen sulphide, of sulphurous 
 acid, and of hydrogen peroxide, on such an acid solution (?). 
 
 c. Add ammonium sulphide, heat and maintain at the boil- 
 ing-point, noting two distinct changes (?), then acidify (?). 
 
 d. Add lead nitrate and barium chloride solutions to sepa- 
 rate portions (?). 
 
 e. Repeat d with potassium dichromate solution (?). Com- 
 pare the results and explain. 
 
 195. Manganates and Permanganates. 
 
 a. Fuse a mixture of 5 g. of potassium hydroxide, 2.5 g. po- 
 tassium chlorate (include only the oxygen, from this, in the equa- 
 
128 CHROMIUM, MANGANESE [ 196 
 
 tion), and 5 g. finely powdered manganese dioxide, at a red 
 heat, in an iron crucible [Storeroom], stirring with the reverse 
 end of a file, until effervescence ceases (?). Add the last ingre- 
 dient gradually. Treat the mass with a small amount of cold 
 water, decant the clear liquid away from the precipitate, and 
 use it in b, c, and d. What is the color of manganate-ion? 
 
 6. Dilute a part of the clear green solution with a very large 
 amount of water in a beaker (?). If no change should occur, 
 pass carbon dioxide into the diluted solution (?). What is the 
 color of permanganate-ion? How does this substance differ 
 from manganate-ion? 
 
 c. Add a few drops of alcohol, and warm (?). 
 
 d. To the rest add a boiling solution of oxalic acid (?). 
 
 e. Repeat c and d with potassium permanganate solution, 
 acidified by adding two or three times its volume of dilute sul- 
 phuric acid (?). 
 
 /. Recall the actions of hydrogen peroxide (60 c), hydrogen 
 sulphide (73 /), and sulphurous acid (83 g) on acidified potas- 
 sium permanganate solution (?). 
 
 196. Reactions of Manganous Salts. Use any manganous 
 salt. What is the color of manganous-ion? 
 
 a. Borax bead in the oxidizing (?) and reducing (?) flames. 
 
 b. Bead of a mixture of sodium carbonate and sodium nitrate 
 on a platinum wire with any manganese compound (?). 
 
 c. To a diluted solution of a manganous salt, add ammo- 
 nium sulphide (?). Is the product a sulphide? 
 
 d. To another portion add sodium hydroxide (?). Divide 
 into two parts. Shake one with air (?). To the other add 
 bromine-water, and warm (?). 
 
 197. Apply to the instructor for two unknown substances. 
 Identify them, and report the result, as directed in 140 and 167. 
 
CHAPTER XXV. 
 
 IRON, COBALT, NICKEL. 
 
 198. Iron. Recall the preparation of iron from an oxide 
 (20), and the action of the metal on dilute acids (16 a) (?). 
 
 199. Reactions of Ferrous and Ferric Salts : I. 
 
 a. Borax bead with any compound of iron (use the oxide) in 
 the reducing (?) and oxidizing (?) flames. 
 
 6. Recall the action of heat upon ferric nitrate (36 a) and 
 upon ferric sulphate (82) (?). 
 
 Prepare a dilute solution of ferrous-ammonium sulphate [Note 
 37, p. 68]. Dilute some ferric chloride solution [Note 40, p. 
 108]. Use portions of these solutions, and add to each the 
 following reagents. Compare the results in each case. 
 
 c. Ammonium hydroxide (?). Shake with air (?). 
 
 d. Potassium ferrocyanide solution (?). 
 
 e. Potassium ferricyanide solution (?) and add much water. 
 /. Ammonium thiocyanate solution (?). 
 
 g. Ascertain by tests whether the commercial hydrochloric 
 acid contains ferrous-ion or ferric-ion (?). Show that it con- 
 tains also sulphuric acid (?). 
 
 200. Reactions of Ferrous and Ferric Salts: II. Reductions 
 and Oxidations. Use diluted ferrous-ammonium sulphate and 
 diluted ferric chloride solutions. 
 
 a. To portions of each solution add ammonium sulphide 
 solution (?). Ascertain in each case whether the action is 
 easily reversible (?). Explain the behavior of the ferric chlor- 
 ide solution [R 756] (?). How could you determine whether 
 the free sulphur was formed before or after the acidification ? 
 
 6. To portions of each solution add potassium iodide solu- 
 tion (?). To a little starch emulsion add a few drops of the 
 ferric-potassium-iodide mixture (?). 
 
 c. Saturate (test?) a portion of the ferrous solution with 
 hydrogen sulphide (?). Explain in terms of the ion-product 
 constant (?). Now add ammonium hydroxide (?). Filter. 
 Wash the precipitate until odorless, and determine whether it 
 is a sulphide or hydroxide (?). Explain in terms of the ion- 
 product constant. 
 
 d. Saturate (test?) a portion of the ferric solution with hy- 
 drogen sulphide (?). Filter. What is the precipitate (test?)? 
 Examine the clear filtrate, and determine, by means of tests 
 
 129 
 
130 IRON, COBALT, NICKEL [201 
 
 from 199 d, e, and /, whether ferric-ion or ferrous-ion is pres- 
 ent (?). Formulate the action of the hydrogen sulphide. 
 
 e. Boil a portion of the ferric solution with excess of pul- 
 verized iron for several minutes. Filter, and apply to part of 
 the clear filtrate (color?) tests selected from 199 d, e, and / (?). 
 Formulate the action of the iron. Use the rest of the filtrate 
 in/. 
 
 /. To 2 c.c. of potassium permanganate solution add a large 
 excess of dilute sulphuric acid (?). Add this mixture drop by 
 drop to the rest of the filtrate from e until the pink color is 
 permanent [R 744] (?). Apply to this liquid tests from 199 d, 
 e, and / (?). Formulate the action in terms of ions alone. 
 Which kinds of ionic chemical change have been illustrated 
 here? How could this action be used for estimating iron? 
 
 What other oxidizing agents convert ferrous into ferric salts 
 (see94/)? 
 
 201. Hydrolysis. Dissolve 0.5 g. each of ferric sulphate and 
 ferrous-ammonium sulphate separately in water, and warm 
 very slightly. Observe the tints by looking downward through 
 the solutions at a piece of white paper (?). Test each solution 
 with Congo red paper, compare (?), and interpret the results. 
 Now add some pure sulphuric acid to each and observe the 
 tints again. Explain. What are the colors of ferrous-ion and 
 ferric-ion, respectively? Which of these allotropic forms of 
 the element is more typically metallic? 
 
 202. Iron- Ammonium Alum. 
 
 a. Weigh 6 g. of ferric sulphate into an evaporating-dish. 
 Weigh out an equimolecular quantity (calculate) of ammo- 
 nium sulphate. Dissolve the salts separately, each in the mini- 
 mum amount of boiling water, mix the solutions, and set 
 aside (?). Describe the crystals (?). Collect them upon a 
 filter, wash them free from the mother-liquor, and dry with 
 filter paper. 
 
 6. Ascertain (see 199 c, d, e) whether iron-alum is a double 
 salt, or a salt of a complex acid (?). 
 
 203. Ferrocyanides and Ferricyanides. Use diluted potas- 
 sium ferrocyanide solution. 
 
 a. Add ammonium hydroxide (?) and compare with 199 c (?). 
 6. Add ammonium sulphide solution and compare with 
 200 a (?). Is ferrous-ion present? 
 
 c. Add bromine-water (shake) in excess, and boil off the 
 superfluous bromine. To a part of the liquid add ferric chloride 
 solution (see 199 d and e) (?). 
 
 d. To the rest of the liquid from c apply the tests in 199 c 
 and / (?). Is ferric-ion present? 
 
206] IRON, COBALT, NICKEL 131 
 
 204. Reactions of Cobalt Salts. For b and c use diluted 
 cobalt chloride solution. 
 
 a. Borax bead in oxidizing (?) and reducing (?) flames. 
 6. Sodium hydroxide solution, first a little (?), then in excess, 
 and warm (?). 
 
 c. Ammonium sulphide solution (?). 
 
 205. Reactions of Nickel Salts. For b and c use diluted 
 nickel sulphate solution. 
 
 a, 6, c. Same as in 204. 
 
 206. Apply to the instructor for two unknown substances. 
 Identify them, and report the result, as directed in 140 and 
 157. 
 
APPENDIX. 
 I. Correction of Barometric Readings. 
 
 To reduce the reading taken at room temperature (Temp.) 
 to the corresponding height of a column of mercury at 0, sub- 
 tract the proper number in the second column (Corr'n) from 
 the actual reading in millimeters. [See Note 31, p. 18.] 
 
 Temp. 
 
 Corr'n. 
 
 Temp. 
 
 Corr'n. 
 
 Temp. 
 
 Corr'n. 
 
 12 
 
 1.6 
 
 17 
 
 2.2 
 
 23 
 
 3.0 
 
 13 
 
 1.7 
 
 18.5 
 
 2.4 
 
 24.5 
 
 3.2 
 
 14 
 
 1.8 
 
 20 
 
 2.6 
 
 25 
 
 3.3 
 
 15 
 
 2.0 
 
 21.5 
 
 2.8 
 
 26 
 
 3.4 
 
 II. Tension of Aqueous Vapor in Millimeters. 
 
 Temp. 
 
 Press. 
 
 Temp. 
 
 Press. 
 
 Temp. 
 
 Press. 
 
 
 
 4.6 
 
 16 
 
 13.5 
 
 26 
 
 25.1 
 
 5 
 
 6.5 
 
 17 
 
 14.4 
 
 27 
 
 26.5 
 
 8 
 
 8.0 
 
 18 
 
 15.4 
 
 28 
 
 28.1 
 
 9 
 
 8.6 
 
 19 
 
 16.3 
 
 29 
 
 29.8 
 
 10 
 
 9.2 
 
 20 
 
 17.4 
 
 30 
 
 31.5 
 
 11 
 
 9.8 
 
 21 
 
 18.5 
 
 31 
 
 33.4 
 
 12 
 
 10.5 
 
 22 
 
 19.7 
 
 32 
 
 35.4 
 
 13 
 
 11.2 
 
 23 
 
 20.9 
 
 33 
 
 37.4 
 
 14 
 
 11.9 
 
 24 
 
 22.2 
 
 34 
 
 39.6 
 
 15 
 
 12.7 
 
 25 
 
 23.6 
 
 100 
 
 760.0 
 
 Aluminium 
 Copper 
 
 in. Specific Heats of Metals. 
 
 0.214 
 0.095 
 
 Iron 
 Lead 
 
 0.114 
 0.031 
 
 Magnesium 
 Zinc 
 
 0.250 
 0.095 
 
APPENDIX 133 
 
 IV. Solubilities of Bases and Salts in Water at 18. 
 
 
 K 
 
 Na 
 
 Li 
 
 Ag 
 
 Tl 
 
 Ba 
 
 Sr 
 
 Ca 
 
 Mg 
 
 Zn 
 
 Pb 
 
 Cl 
 
 32.95 
 3.9 
 
 35.86 
 5.42 
 
 77.79 
 13.3 
 
 o.o 3 i6 
 
 0.0 4 10 
 
 0.3 
 0.013 
 
 37.24 
 1.7 
 
 51.09 
 3.0 
 
 73.19 
 5.4 
 
 55.81 
 5.1 
 
 203.9 
 9.2 
 
 1.49 
 0.05 
 
 Br 
 
 /y- Qt* 
 OO.OU 
 
 4.6 
 
 88.76 
 6.9 
 
 168.7 
 12.6 
 
 0.0 4 1 
 0.0 6 6 
 
 0.04 
 0.0 2 15 
 
 103.6 
 2.9 
 
 96.52 
 3.4 
 
 143.3 
 5.2 
 
 103.1 
 4.6 
 
 478.2 
 9.8 
 
 0.598 
 0.02 
 
 I 
 
 137.5 
 6.0 
 
 177*9 
 84 
 
 161.5 
 8.5 
 
 0.0 6 35 
 0.0 7 1 
 
 0.006 
 0.0 3 17 
 
 201.4 
 3.8 
 
 169.2 
 3.9 
 
 200 
 
 4.8 
 
 148.2 
 4.1 
 
 419 
 6.9 
 
 0.08 
 0.0,2 
 
 F 
 
 92.56 
 12.4 
 
 4.44 
 1.06 
 
 0.27 
 0.11 
 
 195.4 
 13.5 
 
 72.05 
 3 
 
 0.16 
 0.0j92 
 
 0.012 
 0.001 
 
 0.0016 
 0.0 3 2 
 
 0.0076 
 0.0 2 14 
 
 0.005 
 0.0 3 5 
 
 0.07 
 0.003 
 
 NO, 
 
 30.34 
 2.6 
 
 83.97 
 7.4 
 
 71.43 
 7.3 
 
 213.4 
 8.4 
 
 8.91 
 0.35 
 
 8.74 
 0.33 
 
 66.27 
 2.7 
 
 121.8 
 5.2 
 
 74.31 
 4.0 
 
 117.8 
 
 4.7 
 
 51.66 
 1.4 
 
 cio s 
 
 6.6 
 0.52 
 
 97.16 
 6.4 
 
 313.4 
 15.3 
 
 12.25 
 0.6 
 
 3.69 
 0.13 
 
 35.42 
 1.1 
 
 174.9 
 4.6 
 
 179.3 
 5.3 
 
 126.4 
 4.7 
 
 183.9 
 5.3 
 
 150.6 
 3.16 
 
 BrO 8 
 
 6.38 
 0.38 
 
 36.67 
 2.2 
 
 152.5 
 8.20 
 
 0.59 
 0.025 
 
 0.30 
 0.009 
 
 0.8 
 0.02 
 
 30.0 
 0.9 
 
 85.17 
 2.3 
 
 42.86 
 1.5 
 
 58.43 
 1.8 
 
 1.3 
 0.03 
 
 I0 3 
 
 7.62 
 0.35 
 
 8.33 
 0.4 
 
 80.43 
 3.84 
 
 0.004 
 0.0 3 14 
 
 0.059 
 0.0 2 16 
 
 0.05 
 0.001 
 
 0.25 
 0.0 2 57 
 
 0.25 
 0.007 
 
 6.87 
 0.26 
 
 0.83 
 0.02 
 
 0.002 
 0.0 4 3 
 
 OH 
 
 142.9 
 18 
 
 116.4 
 21. 
 
 12.04 
 5.0 
 
 0.01 
 0.001 
 
 40.04 
 1.76 
 
 3.7 
 0.22 
 
 0.77 
 0.063 
 
 0.17 
 0.02 
 
 0.001 
 0.0 3 2 
 
 0.0,5 
 0.0 4 5 
 
 0.01 
 
 0.0 3 4 
 
 80 4 
 
 11.11 
 0.62 
 
 16.83 
 1.15 
 
 35.64 
 2.8 
 
 0.55 
 0.020 
 
 4.74 
 0.09 
 
 0.0 3 23 
 0.0 4 10 
 
 0.011 
 0.0 3 6 
 
 0.20 
 0.015 
 
 35.43 
 
 2.8 
 
 53.12 
 3.1 
 
 0.0041 
 0.0 3 13 
 
 Cr0 4 
 
 63.1 
 2.7 
 
 61.21 
 3.30 
 
 111.6 
 6.5 
 
 0.0025 
 0.0 3 15 
 
 0.006 
 0.0 3 1 
 
 0.0 3 38 
 0.0 4 15 
 
 0.12 
 0.006 
 
 0.4 
 0.03 
 
 73.0 
 4.3 
 
 . . . 
 
 0.0 4 2 
 0.0j5 
 
 C a 4 
 
 30.27 
 1.6 
 
 3.34 
 0.24 
 
 7.22 
 0.69 
 
 0.0035 
 0.0 3 2 
 
 1.48 
 0.030 
 
 0.0086 
 0.0 S 38 
 
 0.0046 
 0.0 3 26 
 
 0.0,56 
 0.0 4 43 
 
 0.03 
 0.0027 
 
 0.0 3 6 
 0.0 4 4 
 
 0.0 3 15 
 0.0 6 5 
 
 C0 3 
 
 108.0 
 5.9 
 
 19.39 
 1.8 
 
 1.3 
 0.17 
 
 0.003 
 0.0 3 1 
 
 4.95 
 0.10 
 
 0.0023 
 0.0 3 11 
 
 0.0011 
 0.0 4 7 
 
 0.0013 
 0.0 3 13 
 
 0.1 
 0.01 
 
 0.004? 
 0.0,3? 
 
 o.o s i 
 
 0.0 4 3 
 
 The upper number in each square gives the number of grams of the 
 anhydrous salt held in solution by 100 c.c. of water. The lower number is 
 the molar solubility, i.e., the number of moles contained in one liter of the 
 saturated solution. * 
 
134 
 
 APPENDIX 
 
 40 - 50 e 
 
 Temperature 
 
 70 89 90 100 
 
APPENDIX 
 
 135 
 
 VI. Degree of lonization of lonogens. 
 
 Except where otherwise specified, the figures give the fraction 
 ionized in a normal, aqueous solution (usually at 18). Sub- 
 traction of the figures from the unity gives the extent to which 
 the ions will unite when brought together in normal concen- 
 tration. At greater dilutions the ionization is greater and the 
 union of ions less. 
 
 ACIDS. 
 
 HN0 3 
 
 0.82 
 
 H.H 2 P0 4 (N/2) 
 
 0.17 
 
 HNO 3 (cone.) 
 
 0.09 
 
 H.HC 2 O 4 (N/10) 
 
 0.50 
 
 HC1 
 
 0.78 
 
 H.HC 4 H 4 6 (N/10 
 
 ) 0.08 
 
 HC1 (cone.) 
 
 0.13 
 
 H.C 2 H 3 2 
 
 0.0 2 4 
 
 HC1 (N/2) 
 
 0.85 
 
 H.C 2 H 3 2 (N/10) 
 
 0.013 
 
 H 2 SO 4 
 
 0.51 
 
 H.HC0 3 (N/10) 
 
 0.0 2 17 
 
 H 2 SO 4 (cone.) 
 
 0.0 2 7 
 
 H.HC0 3 (N/25) 
 
 0.0 2 21 
 
 HBr (N/2) 
 
 0.90 
 
 H.HS (N/10) 
 
 0.0 3 7 
 
 HI (N/2) 
 
 0.90 
 
 H.H 2 BO 3 (N/10) 
 
 0.0 3 1 
 
 HC10 3 (N/2) 
 HMnO 4 (N/2) 
 
 0.88 
 0.93 
 
 HNC (N/10) 
 
 0.0,1 
 
 
 BASES. 
 
 
 KOH 
 
 0.77 
 
 Sr(OH) a (N/64) 
 
 0.93 
 
 NaOH 
 Ba(OH) 2 
 
 0.73 
 
 0.69 
 
 Ba(OH) 2 (N/64) 
 AgOH (N/1783) 
 
 0.92 
 0.39 
 
 NH 4 OH 
 
 0.0 2 4 
 
 HOH 
 
 0.0.1 
 
 Ca(OH) 2 (N/64) 
 
 0.90 
 
 
 
 SALTS. 
 
 
 KC1 
 
 0.75 
 
 Na 2 CO 3 
 
 0.40 
 
 KBr (N/32) 
 
 0.92 
 
 Na.HC0 3 
 
 0.52 
 
 KC10 3 (N/2) 
 
 0.79 
 
 Na,.HP0 4 (N/32) 
 
 0.83 
 
 KN0 3 
 
 0.64 
 
 NaC 2 H 3 2 
 
 0.53 
 
 K 2 S0 4 
 
 0.53 
 
 Na 2 C 4 H 4 O 6 (N/32) 
 
 0.78 
 
 K 2 C0 3 
 
 0.49 
 
 BaCl 2 
 
 0.57 
 
 KMnO 4 (N/32) 
 
 0.92 
 
 CaS0 4 (N/100) 
 
 0.63 
 
 K 2 Cr 2 O 7 (N/32) 
 
 0.94 
 
 CuSO 4 
 
 0.22 
 
 NH 4 C1 
 
 0.74 
 
 AgN0 3 
 
 0.58 
 
 NaCl 
 
 0.67 
 
 CdS0 4 
 
 0.22 
 
 Na 2 S0 4 
 
 0.44 
 
 ZnSO 4 
 
 0.24 
 
 Na 2 SO 8 (N/32) 
 
 0.82 
 
 HgCl, 
 
 (<0.01) 
 
136 APPENDIX 
 
 VII. Electromotive Series. 
 
 The electromotive force of a cell, in which each of the fol- 
 lowing metals constitutes in turn the negative pole (and gold, 
 e.g., the positive), diminishes in the order given. The tendency 
 to enter the ionic condition in a solution already containing 
 the same ion in normal concentration diminishes in the same 
 order, and hence the ionic form of each of these metals (in 
 normal concentration) is discharged and the metal liberated 
 by every metal preceding it in the series. 
 
 Potassium Cadmiurn Arsenic 
 
 Sodium Iron (Fe**) Bismuth 
 
 Barium Thallium Antimony 
 
 Strontium Cobalt Mercury (Hg**) 
 
 Calcium Nickel Silver 
 
 Magnesium Tin (Sn") Palladium 
 
 Aluminium Lead Platinum 
 
 Manganese Hydrogen Gold 
 
 Zinc Copper (Cu") 
 
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X 
 DO 
 
 REFERENCES 
 
 TO THE 
 
 GENERAL CHEMISTRY FOR COLLEGES 
 
 The references given in this Laboratory Outline are to the In- 
 troduction to General Inorganic Chemistry. These references are 
 arranged below in numerical order, and opposite to them are placed 
 the corresponding pages in the General Chemistry for Colleges. 
 
 Inorganic 
 
 College 
 
 Inorganic. 
 
 College 
 
 Inorganic 
 
 College 
 
 Inorganic 
 
 College 
 
 46 
 
 31 
 
 260 
 
 
 387 
 
 263 
 
 600 
 
 398 
 
 49 
 
 32 
 
 264 
 
 187 
 
 390 
 
 265 
 
 601 
 
 398 
 
 57 
 
 41 
 
 267 
 
 191 
 
 440 
 
 293 
 
 601 
 
 399 
 
 63 
 
 67 
 
 268 
 
 191 
 
 448 
 
 299 
 
 607 
 
 403 
 
 70 
 
 51 
 
 269 
 
 192 
 
 459 
 
 306 
 
 621 
 
 413 
 
 71 
 
 51 
 
 273 
 
 
 461 
 
 307 
 
 622 
 
 414 
 
 73 
 
 53 
 
 275 
 
 196 
 
 464 
 
 308 
 
 623 
 
 414 
 
 93 
 
 63 
 
 276 
 
 197 
 
 465 
 
 309 
 
 624 
 
 415 
 
 96 
 
 66 
 
 277 
 
 198 
 
 468 
 
 313 
 
 625 
 
 415 
 
 98 
 
 67 
 
 281 
 
 201 
 
 469 
 
 313 
 
 625 
 
 416 
 
 100 
 
 390 
 
 289 
 
 205 
 
 475 
 
 318 
 
 643 
 
 429 
 
 108 
 
 74 
 
 292 
 
 206 
 
 476 
 
 318 
 
 645 
 
 430 
 
 111 
 
 
 293 
 
 206 
 
 480 
 
 322 
 
 648 
 
 432 
 
 119 
 
 82 
 
 297 
 
 223 
 
 481 
 
 324 
 
 648 
 
 433 
 
 121 
 
 83 
 
 297 
 
 234 
 
 482 
 
 324 
 
 650 
 
 399 
 
 123 
 
 84 
 
 305 
 
 212 
 
 498 
 
 331 
 
 655 
 
 437 
 
 138 
 
 94 
 
 306 
 
 212 
 
 499 
 
 331 
 
 656 
 
 437 
 
 148 
 
 99 
 
 308 
 
 212 
 
 500 
 
 333 
 
 657 
 
 437 
 
 149 
 
 99 
 
 313 
 
 215 
 
 505 
 
 334 
 
 657 
 
 438 
 
 153 
 
 103 
 
 326 
 
 224 
 
 507 
 
 335 
 
 659 
 
 438 
 
 158 
 
 106 
 
 328 
 
 226 
 
 513 
 
 341 
 
 659 
 
 439 
 
 160 
 
 , 106 
 
 332 
 
 229 
 
 518 
 
 343 
 
 660 
 
 439 
 
 161 
 
 107 
 
 334 
 
 224 
 
 521 
 
 345 
 
 673 
 
 
 162 
 
 
 335 
 
 234 
 
 528 
 
 350 
 
 685 
 
 446 
 
 163 
 
 204 
 
 335 
 
 235 
 
 533 
 
 353 
 
 689 
 
 449 
 
 164 
 
 204 
 
 344 
 
 353 
 
 535 
 
 353 
 
 697 
 
 455 
 
 172 
 
 112 
 
 344 
 
 354 
 
 536 
 
 353 
 
 701 
 
 458 
 
 176 
 
 115 
 
 346 
 
 264 
 
 541 
 
 357 
 
 705 
 
 460 
 
 179 
 
 118 
 
 347 
 
 252 
 
 544 
 
 
 710 
 
 464 
 
 180 
 
 118 
 
 349 
 
 238 
 
 553 
 
 
 712 
 
 665 
 
 182 
 
 120 
 
 356 
 
 240 
 
 558 
 
 367 
 
 715 
 
 468 
 
 211 
 
 135 
 
 357 
 
 
 565 
 
 
 715 
 
 469 
 
 230 
 
 161 
 
 359 
 
 231 
 
 571 
 
 374 
 
 718 
 
 470 
 
 232 
 
 
 360 
 
 236 
 
 578 
 
 382 
 
 729 
 
 480 
 
 237 
 
 
 360 
 
 243 
 
 581 
 
 385 
 
 730 
 
 481 
 
 238 
 
 167 
 
 361 
 
 244 
 
 587 
 
 394 
 
 744 
 
 492 
 
 239 
 
 167 
 
 362 
 
 244 
 
 594 
 
 335 
 
 754 
 
 J231 
 
 241 
 
 168 
 
 369 
 
 249 
 
 595 
 
 392 
 
 756 
 
 501 
 
 Hta 
 
 -Jin 
 
 374 
 
 252 
 
 598 
 
 397 
 
 759 
 
 503 
 
 ^^_ 
 
 375 
 
 253 
 
 599 
 
 397