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 GIFT OF 
 Harry East Miller 
 
A LABORATORY MANUAL 
 
 OF 
 
 General Chemistry 
 
 For Use in Colleges 
 
 BY 
 
 WILLIAM C M BRAY, 
 
 Professor of Chemistry in the 
 University of California 
 
 AND 
 
 WENDELL M. LATIMER 
 
 Instructor of Chemistry in the 
 University of California 
 
A LABORATORY MANUAL 
 
 OF | 
 
 General Chemistry 
 
 For Use in Colleges 
 
 BY 
 WILLIAM C. BRAY, 
 
 Professor of Chemistry in the 
 University of California 
 
 AND 
 
 WENDELL M. LATIMER 
 
 Instructor of Chemistry in the 
 University of California 
 
 Copyright, 1921, By William C. Bray and Wendell M. Latimer 
 Price, Fifty Cents 
 
 LEDERER, STREET & ZEUS Co., Publishers 
 BERKELEY, CALIFORNIA 
 
 1921 
 

 PREFACE 
 
 The present laboratory manual has been prepared primarily for the use of 
 students in general inorganic chemistry in the University of California. These 
 students have usually had a year in elementary chemistry in high school, and 
 many of them will take no further work in chemistry. No distinction is made 
 between students on the basis of the various curricula which they are following, 
 as we believe that a course in the fundamentals of general chemistry is equally 
 suitable for all students. 
 
 The laboratory and lecture work are correlated as closely as possible. In the 
 present manual page references are given to Professor Joel H. Hildebrand's 
 "Principles of Chemistry/' MacMillan, 1918, the reference book written for the 
 course. 
 
 The laboratory work is a study of chemical principles, rather than a presen- 
 tation of descriptive material. It is hoped that the division of the manual into 
 Sections, and the statements in the first paragraphs of the various Assignments, 
 will materially assist the student in recognizing the relation between the 
 experimental details and the principles involved. 
 
 When the course extends over two terms, as at the University of California, a 
 satisfactory division is to take Sections I to III in the first term, though in some 
 cases it may be possible also to begin the first Assignment on Qualitative Analysis. 
 It is recommended that the Assignments in the last two Sections be taken in the 
 order noted in the text. 
 
 The following editions of the manual have been printed : Laboratory Directions 
 in Chemistry 1A, edited by William C. Bray, 1915; 21 Assignments. A Labora- 
 tory Manual of General Chemistry, William C. Bray and Ludwig Rosen stein, 
 1916; 26 Assignments. The, 6 s^rne, . revised, by. William C. Bray, 1917; 31 
 Assignments; reprinted 1918, \9i9,l92^J frfiS present manual contains 5 Sections 
 with a total of 30 Assignments, and is an* almost complete revision of the 1917 
 
 manual. A : .::: : ".i :Y ': : :/.'"": 
 
 In the development of this manual from 1912 to the present time a great deal 
 has been contributed by the instructors in the course. We wish especially to 
 acknowledge our indebtedness to Professors G. N. Lewis, J. H. Hildebrand, 
 Edward Booth and E. D. Eastman and to Doctors Ludwig Rosenstein and 
 W. L. Argo. 
 
 WILLIAM C. BRAY, 
 WENDELL M. LATIMER. 
 
 June, 1921. 
 
 * 
 
 a 
 
TABLE OF CONTENTS 
 
 Pago 
 Note to Students 5 
 
 Special Laboratory Directions 5 
 
 List of Apparatus 6 
 
 SECTION I. WEIGHT RELATIONS IN 
 CHEMICAL REACTIONS. 
 
 Assignment 1. A Chemical Reaction: The Synthesis of a Sulfide 
 
 of Copper 7 
 
 Notes on Glass Manipulation 8 
 
 Assignment 2. The Relation between the Mass and Volume of 
 
 Gases: The Determination of the Volume of a Mol of Oxygen.. 9 
 
 Assignment 3. The Reaction between Certain Metals and Hydrochloric 
 
 Acid 11 
 
 Assignment 4. The Analysis of Copper Oxide 13 
 
 Assignment 5. The Reaction between an Acid and a Base in Solution. 
 
 Concentration in Terms of Mols per Liter 15 
 
 Assignment 6. The Titration of Solutions of Acids and Bases : An 
 
 Illustration of Volumetric Analysis 17 
 
 Assignment 7. Volumetric Analysis, Continued : The Determination of 
 
 the Equivalent Weight of an Unknown Acid 19 
 
 SECTION II. IONIC THEORY. 
 
 RAPID REVERSIBLE REACTIONS AND EQUILIBRIUM 
 Assignment 21. Solutions of Strong Electrolytes 21 
 
 Assignment 22. Strong and Weak Acids. The Use of Indicators 
 
 to Measure Hydrogen Ion Concentration 23 
 
 Assignment 23. Strong and Weak Bases. The Use of Indicators 
 
 to Measure Hydroxide Ion Concentration 26 
 
 Assignment 24. Rapid Reversible Reactions and Equilibrium 27 
 
 Assignment 25. The Reversibility of Neutralization Reactions. Hydrolysis.... 30 
 
 SECTION III. REACTIONS OF IONS 
 
 Assignment 31. The Properties of Sodium, Potassium and Ammonium 
 
 Ions. Tests for Chloride, Sulfate and Nitrate Ions 33 
 
 Assignment 32. Calcium Ion 36 
 
 Assignment 33. Carbonate Ion, Bicarbonate Ion and Carbonic Acid 38 
 
 Assignment 34. Sulfates, Chlorides and Nitrates of Copper, Silver and 
 
 Zinc 42 
 
 Assignment 35. Hydroxides of Copper, Silver and Zinc 44 
 
 Assignment 36. Complex Ions of Copper, Silver and Zinc with Ammonia 46 
 
Page 
 Assignment 37. Carbonates and Sulfides of Copper, Silver and Zinc 48 
 
 Assignment 38. Review of the Chemistry of Positive Ions Already 
 
 Considered - 50 
 
 SECTION IV. REACTIONS OF IONS, CONTINUED 
 
 Assignment 41. Oxidation and Reduction. Replacement Reactions. 
 
 Electrical Cells 52 
 
 Assignment 42. Oxidation of Metals to their Ions. Table of Oxidizing 
 
 and Reducing Agents 54 
 
 Assignment 43. Ferrous and Ferric Ions 56 
 
 Assignment 44. Mercurous and Mercuric Ions 58 
 
 Assignment 45. Lead Ion, Chromate Ion 61 
 
 Assignment 46. Stannous and Stannic Ions. Amphoteric Sulfides 61 
 
 Assignment 47. Ions of Aluminum and of Chromium. Peroxides 62 
 
 SECTION V. QUALITATIVE ANALYSIS 
 
 Assignment 51. The Development of a Scheme of Analysis for a 
 
 Limited Number of Positive Ions 65 
 
 (To follow Assignment 38) 
 
 Assignment 52. The Standard Scheme of Analysis. Methods of 
 
 Dissolving Difficulty Soluble Substances 67 
 
 (To follow Assignment 44) 
 
 Assignment 53. Lead, Tin, Aluminum, Chromium and Barium in the 
 
 Scheme of Analysis 70 
 
 (To follow Assignment 47) 
 
NOTE TO STUDENTS 
 
 1. Decide what is the real purpose of each Assignment. Before beginning 
 the experimental work and preferably before coming to the laboratory, read the 
 first paragraphs of the Assignment, study carefully the References, and review 
 earlier, related work. In general look for the connection between the lectures 
 and laboratory work and between each Assignment and the preceding ones. 
 
 2. Master each idea before proceeding to the next one. Form the habit 
 of at once consulting the instructor whenever you are not certain of the cor- 
 rectness of your answer to questions and of your conclusions from the experi- 
 ments. The frequent short written examinations will be of great assistance to 
 you in deciding whether or not you have really understood the work. 
 
 3. An average student who has understood the earlier work can complete 
 an Assignment in the regular laboratory time allotted to it. Students who can- 
 not finish in the stated time, announced by the instructor, must consider them- 
 selves behind the class, and should plan immediately to do extra work at home 
 and in the laboratory. For their convenience the building is open from 
 8:00 A.M. to 4:30 P.M. (Saturdays 8:00 to 12:00). 
 
 4. Your success depends upon your own efforts. If you are in serious 
 difficulty then there is something wrong with your methods. The instructor 
 can assist you in finding out what is wrong, but he cannot do your work for 
 you. No effort on his part can make up for lack of initiative on your part, 
 failure to assume the responsibility of mastering each idea, or inability to 
 improve your method of doing the work. 
 
 SPECIAL LABORATORY DIRECTIONS 
 
 5. A laboratory note book about 6^ inches wide, opening at the side, and 
 not loose-leaved, is recommended. This book and the laboratory manual will be 
 needed at each meeting of the laboratory section, including the first one. 
 
 6. The recording of experiments, observations and conclusions at once in 
 the note book is an essential part of the laboratory work. Entries made from 
 memory or from memoranda on scraps of paper are not records of the experi- 
 mental work. The value of the original record is improved: by dating each 
 day's notes; by numbering the pages of the note book; by never erasing an 
 entry or tearing out a page; by leaving space for additions and corrections; 
 by adopting a plan of distinguishing between the record of the experiments 
 actually performed, and the other entries such as answers to questions, solu- 
 tions of problems, etc. ; by making all calculations neatly at the bottom or side 
 of the page ; and by writing entries in such a way that they will be easily under- 
 stood when the work is reviewed. The descriptions of the experiments per- 
 formed should be very brief when detailed directions are given in the manual, 
 but must be complete when, as in the later work, experiments are devised by 
 the student. A passing grade in the laboratory work will not be given unless 
 the experiments have been completed and the results properly recorded in the 
 note book. 
 
 7. The laboratory desk must be kept neat and dry. An old towel should be 
 used for cleaning the desk top and another towel should be kept clean for use 
 on apparatus. When cleaning apparatus use tap water and a brush to remove 
 all visible dirt and rinse finally with a little distilled water. Before leaving the 
 laboratory the apparatus should be locked up in the desk. 
 
 8. The wash-bottle should only be used to hold distilled water. Before using 
 sterilize the mouth-piece by boiling in water, and never lend or borrow a 
 wash-bottle. 
 
9. The contamination of laboratory reagents can be avoided by keeping each 
 stopper clean and replacing it at once in the proper bottle, and by never pouring 
 anything back into a reagent bottle. 
 
 10. Experiments which give rise to disagreeable or dangerous fumes must 
 always be performed out of doors or in a fume-closet. 
 
 11. At the first meeting of each laboratory section the instructor will dis- 
 tribute the desk keys together with lists of apparatus similar to the one given 
 below. Check the apparatus in the locker, exchange damaged articles at the 
 office, sign the list of apparatus (surname first) and return it to the instructor. 
 Begin work on Assignment I. 
 
 LIST OF APPARATUS 
 
 1. Regular equipment of each locker. Additional articles may be obtained 
 at the office by filling out an "order slip" and signing your name and locker 
 number. Whenever any article is returned to the office sign a "return slip." At 
 the end of the term the locker must contain the same amount of apparatus, no 
 more and no less; the locker must be clean; the apparatus must be clean and 
 dry, and in good condition; glass stoppers must fit, and be protected by the 
 insertion of a piece of paper. 
 
 1 Key. 
 
 5 Beakers, 100 cc., 150 cc., 250 cc., 
 
 400 cc., 600 cc. 
 5 Reagent Bottles. 
 
 2 Sample Bottles, 50 cc. 
 
 1 Graduated Cylinder, 50 cc. 
 4 Flasks, 500 cc., 250 cc., and two 
 125 cc. 
 
 1 Wash-bottle, equipped with glass 
 
 tubing and rubber stopper. 
 
 2 Funnels. 
 
 2 Blue Glasses. 
 
 2 Glass Rods, 12 cm. and 18 cm. 
 30 cm. Glass tubing.* 
 12 Test-tubes. 
 
 1 Watch Glass. 
 
 1 Casserole. 
 
 1 Crucible, with cover. 
 
 2 Evaporating Dishes. 
 1 Crucible Tongs. 
 
 1 Bunsen Burner, with rubber tubing. 
 
 1 Iron Wire.* 
 
 1 Wire Gauze.* 
 
 1 Triangle. 
 
 1 Test-tube Brush.* 
 
 1 Test-tube Holder. 
 
 1 Test-tube Rack. 
 
 1 Package Filter Paper.* 
 
 1 Rule. 
 
 2 Towels. 
 
 Litmus Paper* in a bottle. 
 
 2. The following additional articles may be obtained at the office : 
 
 (a) By signing the regular order slips. Small short-stemmed funnels ; glass 
 flasks, 50 cc. ; matches * ; corks ; rubber stoppers. 
 
 (b) By signing "temporary order ships." Special apparatus for Assignments 
 2 and 4; burettes, with clamps and pinch-cocks; graduated cylinders, 10 cc. 
 and 250 cc. ; thermometers ; paraffin. These articles should be returned 
 when possible during the same laboratory period. 
 
 *Not returnable. At the end of the first term students should retain these articles for 
 use in the second term. 
 
SECTION i v /.\ \ fj \ \- v K - / 
 
 WEIGHT RELATIONS IN CHEMICAL REACTIONS 
 
 ASSIGNMENT 1 
 
 A CHEMICAL REACTION : THE SYNTHESIS OF A 
 SULPHIDE OF COPPER 
 
 References. Hildebrand, Principles of Chemistry, Chapter I, and pages 40-43. 
 
 1. In Assignment 1 we shall study quantitatively a chemical reaction in which 
 two elements, * a metal and a non-metal unite to form a pure compound. The 
 experiment consists in determining the weight of the compound that is formed 
 from a weighed amount of copper when heated with excess of sulfur. From 
 these experimental data, and the atomic weights of the two elements, the relative 
 number of atoms of copper and sulfur in the compound will be calculated. 
 Questions. If 3.04 grams of a certain metal, when burned in oxygen, yield 5.04 g. 
 of a pure compound of the metal and oxygen, what weight of oxygen will combine 
 with 1.00 g. of this metal? What additional information is necessary before the 
 relative number of atoms of the two elements in the compound can be calculated ? 
 
 2. Experiment. Support a clean porcelain crucible, with a cover, on a triangle 
 and heat with the colorless flame of a bunsen burner to low r redness. Let the 
 crucible cool about 10 minutes, and weigh it, with the cover, to 10 milligrams. 
 Note. Do not make any weighings until instructions in the use of the balance 
 have been given. 
 
 3. While the crucible is cooling obtain from the shelf a clean piece of copper 
 wire, weighing about 1 gram, and weigh it to 10 mg. 
 
 4. Place the copper in the weighed crucible and add enough powdered sulfur 
 to cover the copper. Place the cover on the crucible and heat gently (with a 
 small flame) until the sulfur ceases to burn at the edges of the cover, but do not 
 remove the cover while the crucible is hot. Then heat more strongly until the 
 bottom of the crucible just becomes dull red. Again allow to cool about 10 
 minutes and weigh. 
 
 5. Carefully remove the cover and note the appearance of the contents of the 
 crucible, but do not touch the substance. If there is any free sulfur on the cover 
 or the wall of the crucible, replace the cover, heat the crucible and cover, and 
 weigh again. Check the accuracy of the final weight by adding a small quantity 
 of sulfur and repeating the experiment ; continue until two consecutive results 
 agree within 10 mg. At the end of the experiment remove the substance formed, 
 break it and describe its properties. Clean the crucible with hot nitric acid in 
 a porcelain dish, wash with distilled water, dry by heating, and check the 
 original weight. 
 
 6. Questions. What conclusions can you draw from each of the following 
 observations: (a) the properties of the product are different from those of 
 either copper or sulfur; (b) the product appears to be homogeneous and its 
 weight is greater than that of the copper used? What additional evidence is 
 necessary to prove that the product is a pure substance and not a solid solution? 
 
 7. Calculations. Summarize your experimental results and make the calcula- 
 tions necessary to complete a table similar to the following: 
 
 * It is suggested that the student form the habit of writing out the meaning of each 
 italicized word in the text and of giving an example whenever possible. 
 
 [7] 
 
(a) Weight of crucible 
 
 (b) Weight of copper 
 
 (c) Weight of crucible and product 
 
 (d) Weight of product 
 
 (e) Difference between (d) and (b) 
 
 First Second Value 
 Weighing Weighing Chosen 
 
 Calculate what the increase in weight (e) would have been if one gram-atom 
 of copper had been used in the experiment, and enter in the table as line (/) ; 
 show this result to your instructor at once. How does this number compare with 
 the atomic weight of sulphur? How many gram atoms of sulfur have combined 
 with one gram atom of copper? What, then, is the simplest formula of the 
 substance formed? What is the corresponding molecular weight? Write the 
 equation for the reaction, and write out in words what this equation means, in 
 terms of (a) atoms and molecules, (b) gram atoms and mols, (c) grams, and 
 (d) pounds. 
 
 8. Problems. (/) In order to determine the effect of a small error in 
 weighing the copper repeat your calculations, Paragraph 7, using for the weight 
 of copper a value 10 mg. greater than your experimental value. What per cent 
 of the weight of copper is 10 mg. ? This would be the percentage error in the 
 weight of copper if a 10 mg. error in weighing had been made. What is the 
 corresponding percentage error in your value for the weight of sulfur that 
 would combine with one gram atom of copper? 
 
 (2) Calculate the percentage composition of the copper sulfide formed : 
 (a) from your experimental data, and (b) from the formula and the atomic 
 weights of copper and sulfur. Compare the results. 
 
 (j) A sulfide of iron contains 53.8% iron. What is the formula? (In solving 
 this problem consider one gram atom of iron and one gram atom of sulfur as 
 the fundamental units for iron and sulfur. Calculate first the weight of sulfur 
 and then the number of gram atoms of sulfur combined with one or more 
 gram atoms of iron.) 
 
 (4) The formulas of hydrogen sulfide and of ferrous sulfide are H L .S and 
 FeS, respectively. What are their molecular weights. What weight of sulfur is 
 contained in one mol of hydrogen sulfide? In one mol of ferrous sulfide? What 
 weight of hydrogen sulfide could be made from one mol of ferrous sulfide? 
 
 NOTES ON GLASS MANIPULATION 
 
 To bend a piece of ordinary glass tubing, hold it with both hands in a 
 fan-shaped gas flame and rotate it slowly between the thumb and fingers until 
 a 2]/2. to 3 inch portion is uniformly heated and is soft enough to be bent to the 
 proper angle. Set it aside to cool ; glass will remain hot enough to burn the 
 hand for some time after it no longer appears to be hot. 
 
 To cut glass tubing, scratch it with a file at the proper place, grasp it firmly 
 on each side of this mark (protecting the hands with a cloth), and bend the 
 tube away from the mark. 
 
 Always remove the sharp edges of freshly cut glass at once with a file, or by 
 heating in a gas flame. 
 
 To draw down a piece of tubing to a capillary, heat a portion about 1 inch 
 long in an ordinary gas flame to a higher temperature than was necessary in 
 bending the tubing. Hold the tube with both hands and rotate it to ensure 
 uniform heating and prevent the hot portion from sagging. Withdraw from the 
 flame and draw apart slowly to obtain a thick-walled capillary. 
 
 [81 
 
ASSIGNMENT 2 
 
 THE RELATION BETWEEN THE MASS AND VOLUME OF GASES: THE 
 DETERMINATION OF THE VOLUME OF A MOL OF OXYGEN 
 
 References. Hildebrand, Chapter II, and pages 52 and 57 
 
 1. It is often necessary to know the volume of a given mass of a substance, or 
 conversely the mass of a given volume. While in the case of a solid or liquid 
 the relation between the mass and the volume must be determined for the 
 particular substance, the problem is simplified when we are dealing with a gas, 
 since a mol of every gas occupies nearly the same volume under similar condi- 
 tions. In this Assignment we shall determine, under definite conditions of 
 temperature and pressure, the volume of a known weight of oxygen and calculate 
 the volume of one mol at standard conditions. Questions. What information 
 is needed before you can calculate the weight of 10 cc. of mercury? A given 
 solution of sodium chloride in water contains 25.0 percent sodium chloride and the 
 density of the solution is 1.19 g. per cc. ; what volume of solution in cc., and in 
 liters, contains 100 grams of sodium chloride? 
 
 2. When solid potassium chlorate is strongly heated it decomposes with the 
 evolution of oxygen, and the loss in weight gives the weight of oxygen evolved. 
 The volume of the oxygen is determined by measuring the volume of water 
 displaced by an equal volume of oxygen. When pure potassium chlorate is used 
 it must be heated in a hard glass (difficulty fusible) test-tube. The potassium 
 chlorate decomposes more readily and at a lower temperature when a small 
 quantity of manganese dioxide is present. The hard-glass test-tube may then 
 be replaced by a heavy-walled test-tube of ordinary easily-fusible glass ; but care 
 must be taken not to heat the latter tube to a higher temperature than is necessary 
 for the reaction. The manganese dioxide is a catalyst in this reaction, and all 
 of it may be recovered after the potassium chlorate has been decomposed into 
 potassium chloride and oxygen. 
 
 3. Two students working together obtain from the office a heavy-walled glass 
 tube, a rubber stopper, rubber tube, pinch-cock, clamp and small tube containing 
 about 5 grams potassium chlorate. The yellow order slip for "special" apparatus 
 for Assignment 2 should be signed by both students. The apparatus should be 
 returned as soon as the experiment is finished. Each student must keep a complete 
 record of the experiment in his notebook. 
 
 4. Experiment. Set up the apparatus according to the accompanying diagram. 
 Since the apparatus must be gas tight, glass tubing and stoppers must be care- 
 fully fitted. If your tubes and stoppers do not fit, exchange them at the office. 
 Do not use the glass tubing of your wash bottle. Directions for bending glass 
 tubing are given on page 8. 
 
 5. Place in the heavy-walled test-tube about 5 grams potassium chlorate. Add 
 about 50 mg. manganese dioxide, estimating the amount by comparison with the 
 
 sample in the laboratory; mix it with the 
 potassium chlorate, by jarring the tube; and 
 wipe off any powder that is on the outside of 
 the tube or on the inside near its mouth. 
 Assemble the apparatus as before. 
 
 6. Heat the tube gently with a moving gas 
 flame, leaving the pinch-cock open on the 
 rubber tube outlet. Moisture will appear on 
 the walls of the test-tube, which shows that 
 the potassium chlorate and manganese dioxide 
 were not perfectly dry. Gradually warm the 
 tube to within about one inch of the 
 
stopper and at the same time heat the potassium chlorate until gas 
 evolution begins and some water passes over into the beaker. Drive out 
 the moisture with the oxygen by carefully heating the walls of the tube, 
 but do not scorch the rubber stopper. When 50 to 100 cc. water have passed into 
 the beaker, allow the apparatus to cool; close the pinch-cock near the outer end 
 of the rubber tube, which should now be filled with water. Disconnect the 
 chlorate tube, and weigh it with the dry material inside to 10 mg. 
 
 7. Immediately replace the tube in its proper position, open the pinch-cock 
 while the end of the delivery tube is under the water, raise the beaker until the 
 surfaces of the water inside and outside the flask are at the same level, close the 
 pinch-cock again, place a dry beaker under the delivery tube, and open the pinch- 
 cock. If the water continues to siphon into the beaker your apparatus is not gas 
 tight and must be rebuilt. Again heat the tube gradually until gas evolution 
 begins, and continue to heat the potassium chlorate carefully and not too 
 strongly until from 250 to 300 cc. water have been forced over into the beaker. 
 Allow the tube to cool, equalize the level of the water in the beaker and the flask, 
 and then close the pinch-cock on the siphon tube. By means of a 250 cc. graduated 
 cylinder measure the amount of water which the oxygen has forced out of the 
 flask. Finally weigh carefully the hard glass tube containing the partially 
 decomposed chlorate. Question. Why is it necessary to cool the test-tube and to 
 have the water in the beaker and flask at the same level before closing the 
 pinch-cock? 
 
 8. Repeat the experiment, Paragraph 7. The three weighings and two volume 
 measurements give two independent sets of experimental data. Compare the 
 results of your two experiments by preparing a table which will show for each 
 experiment : the weight of oxygen, the corresponding volume of water displaced, 
 the data referred to in 10 to 11 below, and the results of the calculations (12). 
 
 9. Clean the test-tube by placing water in it and shaking. The manganese 
 dioxide is difficultly soluble in water, while both potassium chlorate and potas- 
 sium chloride dissolve readily. Suggest a method of recovering the manganese 
 dioxide and obtaining a mixture of dry potassium chloride and potassium 
 chlorate practically free from manganese dioxide. 
 
 10. To make the calculations it will be necessary to know the barometric 
 pressure at the time you perform each experiment, and the temperature of the 
 water in the flask. The temperature of the water may be assumed to be that of 
 the room, and the barometric pressure will be posted on the blackboard. Enter 
 these data in your notebook before leaving the laboratory. Below is given a table 
 of the vapor-pressure of water at different temperatures : 
 
 Vapor Pressure of Water. 
 
 Temp. C. Vapor Pressure Temp. C. Vapor Pressure 
 
 14 1.2 cm. mercury 24 2.2 cm. mercury 
 
 16 1.3 " ' " 26 2.5 " 
 
 18 1.5 " " 28 2.8 " 
 
 20 1.7 " " 30 3.2 " 
 
 22 2.0 " 32 3.5 " 
 
 11. Questions. Assuming that levels of the water in the beaker and the flask 
 were the same when the pinch-cock was closed, what was the total pressure of 
 the gas in the flask? What was the partial pressure of the water- vapor? Of 
 the oxygen? 
 
 12. Calculations. From each of your two sets of experimental data, by means 
 of the Gas Laws, calculate the volume of 1 mol (32 grams) of pure oxygen at 
 1 atmosphere pressure and O C. Compare your results with the value given in 
 your text book. Your results should not differ from this value by more than 
 
 F 101 
 
5 percent (check your calculations). Show your tabulated results to your 
 instructor, and repeat the experiment if necessary. 
 
 13. Questions. What is the formula of the oxygen molecule? What value is 
 usually accepted as a close approximation for the volume of 1 mol of gas under 
 standard conditions? Use this value to calculate (a) the weight of a liter of 
 hydrogen chloride gas, HC1, under standard conditions; (b) the molecular 
 weight of a gas whose density at standard conditions is known to be 0.001977 
 g. per cc. 
 
 14. Write an equation to represent the decomposition of potassium chlorate, 
 KC1O 3 , into potassium chloride, KC1, and oxygen; and write out what this 
 equation means in terms of (a) molecules; (b) mols, and (c) grams of the 
 substance involved. 
 
 15. Problems, (i) One gram of potassium chlorate is completely decomposed 
 into potassium chloride and oxygen. Calculate (a) the weight of oxygen that 
 could be obtained ; (b) the volume of the oxygen (in liters and in cc.) at standard 
 conditions, and (c) the volume of the oxygen at 27 C. and 750 mm. mercury 
 pressure. 
 
 (2) It is an experimental fact that 2 volumes of carbon monoxide gas react 
 with 1 volume of oxygen to form 2 volumes of carbon dioxide gas. Give the 
 reasoning by which, from this result you can conclude that the molecule of 
 oxygen contains an even number of atoms. 
 
 (5) The formulas of carbon monoxide and carbon dioxide are CO and CO 2 , 
 respectively. Write the equation for the reaction considered in the preceding 
 question, and interpret it in terms of (a) mols, (b) liters, and (c) grams. 
 
 ASSIGNMENT 3 
 THE REACTION BETWEEN CERTAIN METALS AND HYDROCHLORIC ACID 
 
 References. Hildebrand, pages 84-86, and 47-50. 
 
 1. Certain metals, aluminum, zinc, magnesium, etc., react with a solution of 
 an acid, with evolution of hydrogen gas and formation of a salt in solution. In 
 this assignment we shall dissolve a definite weight of a metal in excess of hydro- 
 chloric acid, measure under definite conditions of temperature and pressure the 
 volume of the hydrogen liberated, and calculate the weight of the metal that 
 would form one gram-atom of hydrogen. From this result and the atomic 
 weight of the metal we can then determine : the number of atoms of hydrogen 
 formed when one atom of the metal reacts with the acid, the number of molecules 
 of acid (HC1) which react with one atom of metal, and the formula of the 
 chloride of the metal formed. Questions. What experimental facts and 
 reasoning have led to the conclusion that the formula of the hydrogen molecule 
 is H ? If the valence of hydrogen in HC1 is + 1 what is the valence of the 
 chlorine in this compound? 
 
 2. The instructor will supply to each student a sample of a metal as an 
 "unknown." Experiment. Take a portion of the metal weighing between 0.4 and 
 0.5 g. Clean it, if necessary. Weigh to 5 mg. 
 
 3. Obtain a small short-stemmed funnel at the office, and select a beaker of 
 such size that the funnel when placed in it can be completely covered with water. 
 Place the weighed metal in the beaker, place the inverted funnel over it, and 
 pour freshly distilled water into the beaker until the funnel is completely covered. 
 Note. Tap water contains a relatively large amount of dissolved air, and should 
 
 [in 
 
not be used in this experiment unless it has been heated to boiling to expell the 
 greater part of dissolved air. 
 
 4. Pour distilled water into a half liter flask until the water completely fills the 
 flask. Moisten a piece of filter paper slightly larger than the mouth of the flask, cover 
 the mouth of the flask with paper, taking care that no bubble of air remains below 
 the paper. Invert the flask (over an empty vessel) and lower it into the beaker 
 in such a manner that the stem of the funnel enters the neck of the flask. If 
 a bubble of air enters the flask repeat this operation. The apparatus now 
 consists of a beaker containing a funnel inverted over the metal, and a flask 
 filled with water and inverted over the funnel. Place this apparatus in a large 
 beaker or other vessel, to prevent the water from overflowing on the desk during 
 the remainder of the experiment. 
 
 5. Insert a thistle tube or long-stemmed funnel into the water so that the lower 
 end touches the bottom of the beaker at the rim of the inverted funnel, and 
 through it pour 25 cc. concentrated hydrochloric acid. If the liquid is not stirred 
 the concentrated acid, which is 1.18 times as dense as water, will remain for some 
 time as a layer at the bottom of the beaker, and the metal will be dissolved 
 rapidly. If all the water in the inverted flask is displaced by the hydrogen you 
 have used too much metal or too small a flask and must begin the experiment 
 over again. 
 
 6. When the metal has all dissolved (except a few dark-colored flakes of 
 impurities of negligible weight), place the apparatus in a large basin of tap 
 water and carefully remove the beaker and funnel without allowing any air to 
 enter the inverted flask. Keep the flask in the water for several minutes in order 
 that it may be at the same temperature as the water. Then raise or lower the 
 flask until the level inside and outside the flask is the same. (What is now the 
 pressure of the gas inside the flask?) While the flask is in this position, cover 
 the mouth of the flask with the palm of the hand, remove the flask from the water 
 and invert it. While the gas is escaping, test to prove that it is hydrogen. 
 
 7. Measure the volume of the gas which was contained in the flask by filling 
 the flask completely with water and observing the volume needed. Record in 
 your notebook the barometric pressure (written on the blackboard) and the 
 temperature of the water in which the flask was immersed. 
 
 8. You now have the weight of metal taken, and the volume, at a definite 
 temperature and pressure, of a corresponding amount of hydrogen saturated 
 with water vapor. What is the partial pressure of the water vapor at the 
 temperature of the experiment? What was the partial pressure of the hydrogen 
 in the flask? 
 
 9. Calculate from these data : 
 
 The volume at standard conditions that the hydrogen would occupy if it 
 were dry. 
 
 The weight of the hydrogen. (Use the molecular weight 2.016 and the volume 
 of 1 mol of gas, Assignment 2.) 
 
 The weight of metal that would have liberated 1 gram-atom of hydrogen. 
 Report this value to your instructor, who will tell you the name of the metal if 
 your result is correct to within about 5%. 
 
 By means of the atomic weight of the metal calculate the number of atoms of 
 hydrogen formed when 1 atom of the metal dissolves in acid. 
 
 10. Questions. How many molecules of HC1 react with 1 atom of the metal ? 
 Assuming that the hydrogen of the acid is replaced by the metal, what is the 
 formula of the chloride formed? What is the valence of the metal in this 
 compound? Write the equation for the reaction. 
 
 11. Problems. (i) From the density of hydrogen at standard conditions, 
 
 [12.] 
 
0.00008987 g. per cc., calculate the actual volume of 1 mol of hydrogen. What 
 percentage error did you make in the calculations in Paragraph 9 by assuming 
 the value 22.40 liters. 
 
 (2) If sulfuric acid, H 2 SO 4 , had been used in the above experiment instead of 
 hydrochloric acid, the same result would have been obtained and the final solution 
 would have contained a sulfate of the metal. Write the equation for the reaction. 
 
 (j) The student will have noted that in the first three Assignments we have 
 assumed a knowledge of atomic weights. The arbitrary choice of the unit 
 0=16.00 should present no difficulty, but it is often not clear why a particular 
 value is chosen for the atomic weight of an element rather than some fraction or 
 multiple of this value, e. g., in the case of chlorine why 35.46 is chosen instead 
 of say 17.73 or 70.92. To illustrate how this choice is made on the basis of 
 the experimentally obtainable quantities, molecular weight and percentage 
 composition, the following data may be used. (The molecular weights given in 
 the second column of the table are the accurate values, and not approximate 
 values such as would be obtained directly from the weight of 22.40 liters of gas 
 reduced to standard conditions.) 
 
 No. Grams Chlo- 
 Substance Molecular Weight % Chlorine rine in i Mol 
 
 Chlorine 70.92 100 
 
 Hydrogen chloride 36.47 97.2 
 
 Chlorine oxide (/) 86.92 81.6 
 
 Chlorine oxide (?) 67.46 52.6 
 
 Phosphorus chloride 137.42 77.4 
 
 Carbon chloride 153.84 92.3 
 
 Calculate the values required for the fourth column of the table. What value 
 would you choose for the atomic weight of chlorine? No compound of chlorine 
 has even been made which contains in 1 mol less than 35.46 grams of chlorine. 
 How many atoms of chlorine are contained in a molecule of each of the six 
 substances listed in the table? 
 
 ASSIGNMENT 4 
 THE ANALYSIS OF COPPER OXIDE 
 
 Reference. Hildebrand, Chapter III. 
 
 1. In this Assignment, as an example of chemical analysis, we shall determine 
 the composition of an oxide of copper. The analysis will be made by heating a 
 weighed portion of the oxide in a current of hydrogen and weighing the metallic 
 copper which remains. The oxygen of the oxide unites with the hydrogen to 
 form steam. As in Assignment 1, we shall determine the formula of the 
 compound by assuming the atomic weights of copper and oxygen. It is to be 
 noted, however, that the results could be used to determine the relative atomic 
 weights of copper and oxygen if the formula of the compound were known. 
 Our experimental data, of course, will not be sufficiently accurate to make worth 
 while the calculation of the atomic weight of copper. 
 
 2. Experiment. Two students may work together ; both should sign the order 
 slip for "special apparatus for Assignment 4," which consists of a thick-walled 
 hard glass test tube with a rubber stopper and glass tubes, a thistle tube, and 
 two-holed rubber stopper, a clamp, and a calcium chloride tube (with 2 rubber 
 stoppers, 2 glass tubes, 2 rubber tubes and 2 short glass rods). The apparatus 
 should be returned as soon as the experiment is finished. 
 
 3. Set up a ''hydrogen generator" by fitting your half liter flask with a thistle 
 tube extending through the rubber stopper nearly to the bottom of the flask, 
 
 [13] 
 
and an outlet tube bent at right angles (see note on glass manipulation). Place 
 in the flask about 10 grams of zinc and cover it with about 100 cc. water. To the 
 outlet tube attach a "drying tube" (containing solid calcium chloride, which has 
 the property of absorbing moisture). Make sure that the apparatus is air-tight 
 and wrap the flask in a towel. 
 
 4. Set up the remainder of the apparatus according to the directions of the 
 instructor. Dry the thick-walled glass test-tube that is to contain the copper 
 oxide by heating it gently. When it is cool weigh it carefully, together with 
 any portion of the apparatus that may come in contact with the copper oxide. 
 Place in the tube about 1 gram of copper oxide, wipe off any particles that are 
 not in the portion of the tube that is to be heated. Weigh again carefully to 
 obtain the weight of copper oxide used. Attach the apparatus to the hydrogen 
 generator, pour about 40 cc. concentrated hydrochloric acid down the thistle tube, 
 and allow the hydrogen to pass through the apparatus until it has swept out 
 the air. (Caution. Do not place a flame near the outlet nor heat the oxide 
 while the apparatus contains a mixture of oxygen and hydrogen. A dangerous 
 explosion might result.) Collect the gas in small test tubes by displacement 
 of water and ignite it. Explain how this test may be used to determine when the 
 hydrogen is no longer mixed with oxygen. 
 
 5. When pure hydrogen is passing over the copper oxide, begin to heat the oxide 
 very gently with a small flame and continue to heat cautiously until all the oxide 
 changes color. If moisture collects in the farther end of the tube, drive it out by 
 heating the tube carefully. Question. Where does this moisture come from? 
 
 6. Allow the tube to cool in the current of hydrogen, and weigh it. If you 
 have time, check this result at once by repeating the heating in the current of 
 hydrogen and then weighing; if not, set the tube aside in order that you may do 
 so if the results of the following calculations are unsatisfactory. 
 
 7. Calculate the number of (a) grams; (b) gram atoms of copper that are 
 combined with 1 gram atom of oxygen. What, then, is the formula of this oxide 
 of copper? Repeat the experiment if your results are inconclusive. 
 
 8. Write the equation for the reduction of copper oxide by hydrogen, and 
 interpret in terms of (a) atoms and molecules; (b) gram atoms and mols, and 
 (c) grams. 
 
 9. Calculate the percents of copper and of oxygen in this oxide of copper 
 (a) from your experimental results; (b) from the formula and the atomic 
 weights of copper and oxygen. 
 
 10. Problems. ( i) The formulas of cuprous oxide and cupric oxide are Cu 2 O 
 and CuO, respectively. Write equations for the reactions between the heated 
 oxides and hydrogen to form copper and steam. What weight of copper would 
 be obtained from one gram of each oxide? What is the percentage composition 
 of each oxide? 
 
 (2) What weight of water could be obtained from 1 gram of cupric oxide? 
 What volume would this amount of water occupy at 1 atmosphere pressure and 
 (a) 4 C; (b) 273 C? What volume of hydrogen at 273 C and 1 atmosphere 
 pressure is required to form this amount of water? Is this the amount that 
 would be used in an experiment similar to the one actually performed? 
 
 (3) What are the formulas of cuprous sulfide, cupric sulfide, and hydrogen 
 sulfide? Cuprous and cupric sulfides are also reduced to copper when they are 
 heated in a current of hydrogen; hydrogen sulfide is formed. Write equations 
 for the reactions. 
 
 14 
 
ASSIGNMENT 5 
 
 THE REACTION BETWEEN AN ACID AND A BASE IN SOLUTION 
 CONCENTRATION IN TERMS OF MOLS PER LITER 
 
 References. Hildebrand, Chapter V, pages 76-80, Chapter VIII, pages 105 
 
 and 106. 
 
 1. In this Assignment we shall study the reaction between sodium hydroxide 
 and hydrochloric acid in solution and shall determine the amount of salt that is 
 formed from a measured volume of sodium hydroxide solution. In the preceding 
 Assignments the amount of a substance was determined by weighing or by 
 measuring the volume of the pure substance. It is often more convenient to 
 determine the quantity of a substance by measuring the volume of a solution 
 which contains a known amount of the substance in a unit volume of the solution. 
 The amount of the substance in a unit volume of solution is called the 
 concentration. Question. If the concentration of a salt solution is known to be 
 10 g. per liter, what volume would you measure out in order to have 0.5 g. of salt? 
 
 2. It is necessary first to examine separately the properties of the three solutions, 
 the base, acid and salt. Experiment. Prepare dilute solutions of sodium 
 hydroxide, NaOH, and of hydrochloric acid, HC1, by diluting 5 cc. of the 
 laboratory solution of each with 50 cc. of distilled water, and also a dilute solution 
 of NaCl by dissolving between one and two grams of the salt in 50 cc. of distilled 
 water. 
 
 To 10 cc. portions of each of the three solutions add a few drops of litmus 
 solution. Repeat using phenolphthalein. 
 
 Taste each solution by dipping a glass rod into the liquid and touching it to 
 the tongue. (Caution. Do not taste any substance in the laboratory unless 
 directed to do so.) 
 
 Test a drop of each solution in a colorless gas flame by means of an iron (or 
 platinum) wire. A yellow flame proves the presence of sodium. 
 
 Evaporate to dryness in a casserole 1 cc. of HC1 solution, and of NaCl solution. 
 In each case examine if there is a residue. Question. What conclusion can you 
 draw in regard to the volatility of water and hydrogen chloride as compared to 
 sodium chloride? What result would you predict if a solution containing both 
 NaCl and HC1 were evaporated? Pure sodium hydroxide is a stable non-volatile 
 substance which would be left as a solid when a solution containing it is 
 evaporated to dryness. (Caution. Do not evaporate alkaline solutions to dryness. 
 Glass and porcelain are slowly attacked by hot concentrated alkaline and a porcelain 
 dish is spoiled if an alkali residue is heated strongly in it.) 
 
 Summarize in a table the properties of the three solutions examined above. 
 
 3. Experiment. Take 40 cc. of your laboratory NaOH solution, place it in a 
 clean half liter flask and dilute with 440 cc. of distilled water. Shake the flask in 
 order that the solution shall be uniform throughout. Cork the flask and label it 
 "NaOH solution for Assignments 5, 6 and 7." Question. What approximately 
 is the ratio of the initial volume of the sodium hydroxide solution to the final 
 volume? 
 
 4. Dry a porcelain dish and watch-glass large enough to cover it. Weigh the 
 dish and watch-glass to 10 mg. Measure out in your graduated cylinder as 
 accurately as possible 50 cc. of the sodium hydroxide solution prepared above. 
 Pour the solution into the weighed evaporating dish, and sufficient phenolph- 
 thalein to give a pink color and then add small portions of your laboratory 
 hydrochloric acid, stirring after each addition, until the solution is colorless. 
 Approximately 5 cc. of the acid will be required to give the colorless solution. 
 If the addition of the last portion of acid is made, a few drops at a time, it will 
 
 f 151 
 
be observed that the color changes abruptly. An excess of 1 cc. of HC1 may 
 now be added. 
 
 5. Place the dish containing the solution on your wire gauze and heat until 
 the solution begins to boil. Reduce the size of the flame and allow the solution 
 to boil gently, or to evaporate slowly just below the boiling point, until the 
 bottom of the dish is covered with solid material. Then, to avoid loss from 
 bumping and spattering during the evaporation to dryness, cover the dish loosely 
 with the watch-glass, leaving an open space at one side, and continue the heating, 
 first with a small flame, then more strongly until no further trace of water-vapor 
 is expelled. Allow the dish and residue to cool for 10 minutes while covered 
 with the watch-glass and again weigh to 10 mg. Heat the dish and residue gently 
 for five minutes longer, let cool, and weigh again. If the two weights are not the 
 same within 20 mg. repeat this process until two weights are obtained which 
 check to 20 mg. 
 
 6. Dissolve a portion of the residue in a small amount of water and test as 
 in Paragraph 2. State what evidence you have that a reaction has taken place 
 between the acid and base, and that the solid residue obtained is sodium chloride. 
 Write the equation for the neutralisation reaction between NaOH and HC1 and 
 interpret it in terms of (a) molecules, (b) mols, (c) grams of the substances 
 involved. 
 
 7. Calculations. From the weight of NaCl found in Paragraph 5, calculate 
 (a) the weight of NaOH which must have been present in the 50 cc. of NaOH 
 solution, (b) the number of grams of NaOH in 1 cc. of the solution, (c) the 
 concentration of the NaOH in grams per liter? As we shall see in the next 
 assignment this value may not be very accurate. In addition to errors such as 
 that made in measuring out the volume of the NaOH solution, the laboratory 
 NaOH contains small amounts of impurities, as NaCl. 
 
 8. The reaction just considered is typical of the reaction between any acid and 
 any base. In every case H of the acid unites with OH of the base to form water. 
 Write the equations for the reactions between the following bases and acids, and 
 interpret each equation in terms of mols of the substances involved : 
 
 Sodium hydroxide and nitric acid 
 
 Sodium hydroxide and sulfuric acid (to form two molecules of water) 
 
 Barium hydroxide and hydrochloric acid 
 
 Barium hydroxide and sulfuric acid. 
 
 9. Since the mol is a convenient unit of weight to use in studying chemical 
 reactions, concentration is often expressed in terms of the number of mols of 
 substances in a liter of solution. {Definition. A solution which contains in one 
 liter one mol of dissolved substance is called a molal solution; one which 
 contains in one liter one-tenth mol of dissolved substance is called a tenth molal 
 solution, etc. Molal hydrochloric acid is designated thus: M HC1; tenth molal 
 sulfuric acid would be written 0.1 M H 2 SO 4 , etc. Questions. Calculate the 
 number of mols of NaCl obtained in Paragraph 5. How many mols of NaOH 
 were there in 50 cc. of solution ? What then is the concentration of your NaOH 
 in mols per liter? If the laboratory solution were exactly 6 M (which probably 
 is not the case) what would this concentration be, as a result of the twelve fold 
 dilution in Paragraph 3? 
 
 10. Problems, (i) How many (a) mols, (b) grams of sulfuric acid are in 
 50 cc. of 0.2 M H 2 SO 4 ? 
 
 (2) What is the concentration in mols per liter of a solution which contains 
 5.8 g of NaCl in 125 cc. of solution. 
 
 16 
 
ASSIGNMENT 6 
 
 TITRATION OF SOLUTIONS OF ACIDS AND BASES : AN ILLUSTRATION 
 OF VOLUMETIC ANALYSIS 
 
 Reference. Hildebrand, Chapter VIII, pages 106-107. 
 
 1. In Assignment 5 we learned that an acid and a base in solution will neutralize 
 each other. In Assignment 6 we shall see how this reaction can be used in 
 determining the concentration of one of these solutions when that of the other is 
 known. The operation is called a titration. It is evident that a pure solution 
 of a salt may be prepared by mixing the corresponding acid and base in exactly 
 the right proportion; this end-point in the titration is determined by means of a 
 suitable indicator .The relative concentration of the two solutions can be calcu- 
 lated when the relative volumes of the two solutions used in the titration are 
 known. Questions. How many mols of NaOH are required to neutralize 
 exactly 0.01 mol of (a) HC1, (b) H,SO 4 ? How many cc. of 0.50 M NaOH 
 solution will exactly neutralize 0.01 mol of (a) HC1, (b) H 2 SO,.? 
 
 2. Experiment. Prepare 300 cc. approximately 0.5 M HC1 from your labora- 
 tory 6 M solution, place it in a flask and shake it. Cork the flask and label it 
 "HC1 solution for Assignments 6 and 7." Clean a small flask and label it 
 "known H 2 SO 4 solution/' rinse the flask with distilled water and set it aside to 
 drain in order to have it ready for use later in this Assignment. 
 
 3. Experiment on the determination of an end-point and the choice of the 
 indicator to be used in the titration. Dissolve approximately 0.5 g. NaCl in about 
 50 cc. water, add 2 drops phenolphthalein and stir the solution. Add NaOH 
 solution (approximately 0.5 M) drop by drop, stirring after each drop is added, 
 and note how many drops are needed to give a distinct color. Then determine 
 how many drops of your 0.5 M HC1 solution are required to decolorize the 
 solution. Note. Small drops are conveniently added from a glass tube drawn 
 out to a point, the solution being held in the tube by placing the finger over the 
 upper end of the tube. Be sure that the tube is clean ; before using, rinse it once 
 with the solution. Repeat the experiment with litmus solution instead of phenol- 
 phthalein, and determine whether the change in color gives a satisfactory end- 
 point for the titration of HC1 and NaOH solutions. Finally determine the 
 nature of the end-point with each indicator when about 1 g. sodium sulfate, 
 Na 2 SO 4 - 10 H 2 O is used instead of NaCl. 
 
 4. The volumes of the solutions used in a titration are measured by means of 
 burettes; these are uniform glass tubes graduated in cubic centimeters and tenths 
 (or fifths) of cubic centimeters. Two burettes will be placed on your desk before 
 the beginning of the Assignment. At the end of the period rinse them with 
 distilled water and leave them on your desk. 
 
 5. Experiment. Fill each burette with distilled water. Air may be removed 
 from the small tube below the pinch-cock by tilting the tip upward and allowing 
 the liquid to flow through the pinch-cock. Question. Why is it necessary to 
 remove any bubble of air trapped in this small tube? Practice reading a burette: 
 bring your eye to the same level as the liquid and note the reading of the burette 
 corresponding to the bottom of the meniscus; repeat until consecutive readings 
 check to better than 0.05 cc. Question. Why is it important to have the eye at 
 the same level as the liquid before making a reading? Never attempt to adjust 
 the volume of the solution in a burette so that the reading will be some exact 
 amount. Allow the water to flow slowly out of the burette. If drops remain on 
 the inner surface of a burette, exchange it at the office for a clean one. 
 
 6. Rinse one burette with a little of your approximately 0.5 molal HC1 solution, 
 and fill the burette with this solution. Rinse and fill the other burette with the 
 0.5 molal NaOH solution. 
 
 [ 171 
 
7. Record the readings of the burettes side by side in your note-book ; run 
 about 15 cc. of the acid solution into a clean beaker or flask standing on white 
 paper, and record the final burette reading under the initial reading. Add two 
 drops of phenolphthalein, and about 20 cc. distilled water. Then run in the 
 sodium hydroxide solution from the other burette, a little at a time, and towards 
 the end, very carefully, a drop or two at a time, stirring the mixture constantly, 
 until the faintest perceptible permanent pink color is obtained. Wash down the 
 inside of the beaker by means of a jet of water from the wash bottle. If two 
 much of the basic solution is added, decolorize the solution by adding a little of 
 the acid and determine the end-point again. Record the final readings of each 
 burette, and the actual volumes of each solution used in titration. Calculate the 
 volume of sodium hydroxide necessary to neutralize one cubic centimeter of the 
 hydrochloric acid. 
 
 8. Repeat this experiment, using about 20 cc. of the acid solution, and in each 
 case make the same calculations. Do not fill up the burette each time unless there 
 is not enough solution in it for the titration. 
 
 9. Questions. If the error in measuring out a volume of solution by means of a 
 burette is 0.10 cc. what is the percentage error if 1 cc. of solution is measured? 
 If 20 cc. of solution are measured? Why is it important not to use less than 
 10 cc. in any titration? 
 
 10. Compare the volume ratios calculated from two titrations. If the result 
 differs from the average by more than 1%, perform additional titrations until 
 you are satisfied that you have determined the volume ratio with an accuracy 
 better than 1%. 
 
 11. Questions, (a) From an examination of the equation for the neutraliza- 
 tion of sodium hydroxide by hydrochloric acid state the ratio of the number of 
 mols of acid and base added to the beaker when exact neutrality was reached. 
 
 (b) From your average volume ratio state which solution, acid or base, is the 
 more concentrated, and what is the ratio of the concentrations. 
 
 12. Take your clean, dry, labelled flask, Paragraph 2, to the office to obtain a 
 sulfuric acid solution of known concentration. 
 
 13. Empty the HC1 out of the burette, rinse it with about 5 cc. of the sulfuric 
 acid solution of known concentration, and fill it with the sulfuric acid 
 solution. Determine the volume ratio as before from three (or more) titrations. 
 From the volume ratio and the reaction between sodium hydroxide and sulfuric 
 acid calculate the concentration of the sodium hydroxide solution in mols per 
 liter. Calculate also the concentration of your hydrochloric acid solution. 
 
 14. Make a list of the sources of error. (Many of them have been mentioned 
 in the above directions.) 
 
 15. Save the remainder of the NaOH and HC1 solutions, whose concentration 
 you have determined, in corked flasks for use in Assignment 7. 
 
 16. Problems, (i) How many cc. of 0.01 M Ba(OH) 2 will be required to 
 neutralize 10 cc. of 0.5 M HC1? 
 
 (2) What is the concentration in mols per liter of a sulfuric acid solution, 
 25 cc. of which neutralizes 20 cc. of 0.20 M NaOH? 
 
 [18 
 
ASSIGNMENT 7 
 
 VOLUMETIC ANALYSIS, CONTINUED: THE DETERMINATION OF THE 
 EQUIVALENT WEIGHT OF AN UNKNOWN ACID 
 
 Reference. Hildebrand, Chapter VIII, pages 108-111. 
 
 1. In Assignment 7 there will be introduced another unit of quantity, the 
 equivalent. This unit and the corresponding unit of concentration, equivalents 
 per liter, are frequently more convenient than the units, mol and mols per liter 
 The mol and equivalent are identical for HC1, HNO 3 , NaOH, NaCl, etc., but a 
 mol of H 2 SO 4 , Ba(OH) 2 or Na 2 SO 4 , etc., contains two equivalents. A solution 
 which contains one equivalent in a liter is called a normal solution and is desig- 
 nated i N. The convenience of this unit of concentration depends upon the fact 
 that when one equivalent of any acid reacts with one equivalent of any base 
 the resulting solution contains one equivalent of the corresponding salt. 
 Question. What is the normal concentration of a molal solution of H 2 SO 4 , 
 NaOH and Na 2 SO 4 respectively? How many equivalents of acid can be 
 neutralized by 10 cc. of 0.1 N NaOH? Calculate the normal concentrations of 
 your solutions of NaOH, HC1 and H 2 SO 4 used in Assignment 6. 
 
 2. Weighed .portions of an unknown solid acid will be titrated with your NaOH 
 solution. By means of the concentration of the NaOH solution (determined in 
 Assignment 6) the number of grams in one equivalent of the acid will be calcu- 
 lated. The correctness of this result depends of course, on the accuracy of your 
 work in Assignment 6. It is to be noted that if you had started with a solid and 
 of known equivalent weight, you could have used the results obtained in this 
 Assignment to determine the concentration of the sodium hydroxide solution. 
 
 3. Experiment. Obtain from the office a sample bottle containing approxi- 
 mately 2 grams of a crystalline acid of unknown composition. Clean two % liter 
 flasks and label them No. 1 and No. 2. Weigh the sample bottle with its cork and 
 contents. Remove the cork, taking care that none of the solid which may be 
 sticking to it drops off, and shake about one gram into flask No. 1. Replace the 
 cork and weigh again. Shake the remainder of the sample into flask No. 2, and 
 weigh the empty sample bottle and -cork. All weighings should be made to 5 mg. 
 
 4. Dissolve the contents of each flask in about 50 cc. of distilled water, add 
 two drops of phenolphthalein and titrate to the appearance of the first pink color 
 with the sodium hydroxide solution prepared in Assignment 8. If you should 
 pass the end-point in a titration add from a second burette an acid solution of 
 known concentration until the pink color is discharged, and again titrate to the 
 end-point with the sodium hydroxide. 
 
 5. For each sample calculate the number of cubic centimeters of NaOH 
 solution needed to neutralize 1 gram of the acid. If the results differ by more 
 than 2% repeat the experiment. 
 
 6. From your two (or more) measurements obtain an average value of the 
 number of cubic centimeters of sodium hydroxide solution per gram of acid. 
 From this average value and the concentration of the NaOH solution calculate 
 the number of mols of NaOH needed to neutralize 1 gram of the acid. 
 
 7. Questions. How many equivalents are in one gram of the acid? How 
 many grams are in one equivalent of the acid? 
 
 Report this value to your instructor at once. If your result is unsatisfactory 
 check your calculations in Assignments 6 and 7 and repeat as much of the work 
 as is necessary. 
 
 8. Problems, (i) Among the acids suitable for this Assignment are: oxalic 
 H 2 C 2 O 4 2H 2 O ; citric, H 3 C 6 H 5 O 7 H 2 O ; tartaric, HX 4 H 4 O e and potassium acid 
 sulfate, KHSO 4 . Write the equations for the reaction between NaOH and (a) 
 
 [19] 
 
oxalic acid to form Na 2 C 2 O 4 , and (b) KHSO 4 ; and calculate the equivalent 
 weight of the acid in each of these reactions. 
 
 (2) What is the normal concentration of a sulfuric acid solution, 25 cc. of 
 which neutralizes 20 cc. of 0.20 AT NaOH? 
 
 (j) Chemically pure ("C. P.") sulfuric acid, nitric acid, hydrochloric acid 
 and ammonia, as supplied by the manufacturers, are concentrated aqueous 
 solutions of these substances. The concentration of each solution is guaranteed 
 not to be less than a certain minimum value, and this is tested by measuring the 
 density (or the specific gravity). The following table contains the density and 
 the percentage composition by weight of the concentrated laboratory reagents. 
 
 Density Concentration, 
 
 g. per cc. % by weight Equivalents per liter 
 
 H 2 SO 4 1.84 95.6% H 2 SO 4 
 
 HNO 3 1.42 69.8% HNO 3 
 
 HC1 ' 1.19 37.2% HC1 
 
 NH 4 OH 0.90 57.5% NH 4 OH 
 
 For each solution calculate the normal concentration and record the results in 
 the fourth column of the table. 
 
 Caution! The concentrated acids, especially sulfuric and nitric, produce 
 dangerous burns and should not be used carelessly. When you are directed to 
 use one of these acids carefully pour just enough of it for your experiment into 
 a dry beaker. 
 
 20 
 
SECTION II 
 
 IONIC THEORY 
 
 RAPID REVERSIBLE REACTIONS AND EQUILIBRIUM 
 
 ASSIGNMENT 21 
 SOLUTIONS OF STRONG ELECTROLYTES. IONIC EQUATIONS 
 
 Reference. Hildebrand, Chapter X, pages 124-137. 
 
 1. This Assignment, which contains no experimental work, is introduced in 
 order that the student may become familiar with the fundamental ideas under- 
 lying the Ionic Theory before proceeding to use these ideas in the following 
 Assignments. 
 
 2. The use of the terms acids and bases in designating two distinct groups of 
 substances implies that the members of each group have a set of properties in 
 common. The properties characteristic of all acid solutions, which were observed 
 in Assignment 5, are ascribed to a substance called hydrogen ion, represented 
 by the symbol H + ; and those of basic solutions to the substance hydroxide ion, 
 written OH~. In addition to the properties of the hydrogen ion, each acid in 
 solution has a group of properties different from those of any other acid but 
 common to solutions of all salts of that acid. Thus, hydrochloric acid has a set 
 of properties which is characteristic of solutions of all chlorides and is ascribed 
 to a substance called chloride ion, Cl~. Likewise, a solution of sodium hydroxide 
 has a group of properties which is characteristic of solutions of all sodium salts 
 and which is attributed to the sodium ion, Na + . Questions. What is the evidence 
 from freezing point data that there are approximately 2 mols of substance pres- 
 ent when one mol of hydrogen chloride is dissolved in water? How does the 
 electrical conductivity of hydrochloric acid solution support the idea that the 
 molecule of hydrogen chloride in solution is broken up into two new substances? 
 In what way does a chloride ion differ from an atom of chlorine, and a hydrogen 
 ion differ from an atom of hydrogen? List briefly differences in properties of 
 the substances hydrogen ion and hydrogen gas. 
 
 3. Many substances in dilute solution may be considered as completely ionized. 
 These substances are called strong electrolytes and include : 
 
 Practically all salts 
 
 A few acids as HC1, HNO 3 and H 2 SO 4 , and 
 
 A few bases as NaOH, KOH and Ba(OH) 2 . 
 
 The student should memorize this list and should form the habit of thinking 
 of solutions of strong electrolytes in terms of the ions present and not merely in 
 terms of the specific solid, liquid or gas used in making the solution. It is impor- 
 tant to realize, however that strong electrolytes are not ionized in the solid or 
 gaseous state ; thus, while hydrochloric acid and sodium chloride solutions consist 
 of H + and Cl~ and Na + and Cl~, respectively, gaseous HC1 and solid NaCl are 
 not ionized, and each has its own specific properties. Questions. What are the 
 principal substances present in each of the following solutions and what is the 
 approximate concentration of each substance in mols per liter : 
 
 (/) A solution which contains 0.1 mol of H 2 SO 4 in 1 liter. (Answer. H + and 
 SO 4 ~" at concentrations 0.2 M and 0.1 M, respectively). 
 
 (-2) A solution which contains 0.2 mol of NaOH in 1 liter. 
 
 (5) A solution which contains 0.1 mol of Na SO 4 and 0.1 mol of NaCl in 
 1 liter. 
 
 [21] 
 
(4) A solution which is made by mixing equal volumes of (i) and (2). 
 
 4. Ionic equations. Having realized what substances are present in solutions 
 of strong electrolytes we are now in a position to interpret reactions involving 
 such solutions. We shall first consider what is the ionic reaction when a strong 
 acid neutralizes a strong base, and shall take as an example the reaction between 
 sulfuric acid and sodium hydroxide solutions. The equation 
 H,SO 4 + 2NaOH = 2H 2 O + Na 2 SO 4 
 
 is a statement of the reaction that takes place whenever H 2 SO 4 and NaOH 
 solutions are mixed, and we have seen that it may be interpreted in terms of 
 molecules, and of mols, equivalents, grams, or any other weight units. It records 
 no. experimental details, such as volumes or concentrations of solutions, or the 
 use of either reagent in excess; and additional notes are necessary when such a 
 record is desired. 
 
 This equation can be interpreted according to the ionic theory (a) by adding 
 a note that H 2 SO 4 , NaOH and Na 2 SO 4 are strong electrolytes and that H 2 O is a 
 weak electrolyte, or (b) by rewriting the equation to give the same information: 
 
 (2H + + SO 4 ~) + (2Na + + 2OH~) = 2H 2 O + 2Na + + SO 4 ~. 
 It is evident that the Na + and SO 4 ~~ shown to be in the final solution were present 
 in the two initial solutions, and that these substances have undergone no change 
 during the reaction. It is incorrect to say that "sodium ion and sulfate ion have 
 combined/' The reaction that has taken place is simply the formation of the weak 
 electrolyte water: 
 
 5. The statement that a substance is a weak electrolyte is equivalent to saying 
 that the ions of that substance cannot exist in the presence of each other except 
 at low concentrations and it is obvious that any reaction which involves the 
 formation of a weak electrolyte from its ions may be expected to take place. 
 Accordingly, we shall be able to predict a number of reactions when we have 
 classified the substances involved as strong or weak electrolytes in solution. 
 
 6. Another reaction which involves the ions of a strong electrolyte is the 
 precipitation of a sparingly soluble salt. Again the statement that a salt, such as 
 silver chloride, AgCl ; is sparingly soluble is equivalent to saying that its ions 
 cannot exist together in solution except at the low concentrations corresponding 
 to the solubility of the salt. Write the ionic equation for the reaction that takes 
 place when dilute solutions containing equivalent amounts of NaCl and AgNO 3 
 are mixed. What substances are present at high concentrations in each of the 
 initial solutions and in the final solution? Write the ionic equation for the 
 reaction. Consider next the case in which silver nitrate solution in excess is 
 added to a sodium chloride solution, answer the same question and write the 
 ionic equation. It should be obvious that the reaction is the same in both cases, 
 and that, therefore, the two equations should be identical. 
 
 7. Does anything happen in the following experiments ? A dilute solution of 
 sodium chloride is mixed with a dilute solution of (i) potassium nitrate, 
 (2) nitric acid? 
 
 8. Problems, (i) Write ionic equations for the following reactions: (a) A 
 precipitate of barium sulfate, BaSO 4 , is formed by mixing solutions of barium 
 chloride and sulfuric acid, (b) A sodium sulfate solution is evaporated to dry- 
 ness. (c) Hydrogen chloride gas is dissolved in water. 
 
 (2) The solubility of lead iodide, PbI 2 , is 0.002 mol per liter at 18. What 
 is the concentration of the ions present in a saturated solution? 
 
 (j) If the freezing point of water is lowered approximately 1.86 per mol of 
 substance in solution in 1000 grams of water, what is the freezing point of a 0.1 
 molal solution of (a-) a non-electrolyte, (b) hydrochloric acid, (c) barium 
 chloride, BaQ 2 ? 
 
 [221 
 
ASSIGNMENT 22 
 
 STRONG AND WEAK ACIDS. THE USE OF INDICATORS TO MEASURE 
 HYDROGEN ION CONCENTRATION 
 
 Reference. Hildebrand, Chapter X, pages 138-142. 
 
 1. Indicators can be used to measure the concentration of hydrogen ion, or of 
 hydroxide ion, in a solution. The color change for each indicator occurs in a 
 definite range of concentrations of hydrogen ion (or hydroxide ion), which is 
 characteristic of the indicator. In the present Assignment, by using solutions of 
 known concentration of the strong acids listed in Assignment 21, we shall develop 
 a method of determining approximately the concentration of hydrogen ion in 
 any acid solution. This indicator method will then be used to measure the concen- 
 tration of hydrogen ion in solutions of a typical weak acid. Two reactions 
 involving this acid will be studied. 
 
 2. Experiment. Prepare 60 cc. N HC1 by adding distilled water to the proper 
 volume of the 6 N laboratory reagent and shaking the mixture. From this 
 solution, or from your known HC1 solution, Assignment 6, prepare between 50 
 and 100 cc. of 0.10 N HC1, and from this solution prepare 50 to 100 cc. of 0.001 
 N HC1. On account of the error in measuring volumes by means of a graduated 
 cylinder, and the probable variation of the concentration of the laboratory 
 solution from 6 N, the concentrations of these solutions are known only 
 approximately. 
 
 3. Pour into marked test tubes 10 cc. of each HC1 solution (N, 0.1 N, 0.01 N, 
 0.001 N) and pour 10 cc. of water into a fifth test-tube. Add to each solution 
 from a glass tube a single drop of methyl violet solution. Hold the tubes in a verti- 
 cal position over a piece of filter paper, look down through the surface, record the 
 color of each solution and note the smallest concentration of hydrochloric acid 
 that shows with this indicator a color different from that of water. If the 
 indicator solution is so dilute that one drop does not produce a distinct color add 
 1 or 2 more drops, but record the number of drops, and use the same number in 
 all the tubes. State how the indicator, methyl violet, may be used to determine 
 the approximate concentration of a hydrochloric acid solution. (The color in the 
 more concentrated solution will fade on standing. It may be restored by adding 
 another drop of the indicator.) Note. Set aside the remainder of the 0.01 N 
 and 0.001 TV" HC1 for use later in this Assignment, Paragraphs 5 and 8. 
 
 4. Experiment. Repeat the experiment with nitric acid or with sulfuric acid, 
 using the same concentrations as before (N, 0.10 N, 0.01 N, 0.001 N). Compare 
 the colors obtained with the different acids. Questions. Are the colors charac- 
 teristic for each acid? If not, what substance determines the color? If 
 hydrochloric acid is completely ionized what conclusion can you draw with respect 
 to the ionization of nitric acid and sulfuric acid? Note. Save the 0.001 N 
 solution for later use, Paragraph 5. 
 
 5. If you have performed the above experiments correctly and understood 
 them you will realize that the indicator methyl violet can be used to determine, 
 approximately, concentrations of hydrogen ion between N and 0.001 N. 
 However, while the concentration of H + in the dilute solution, 0.001 N, (which is 
 often written 10~ 3 N), is small compared with the normal solution, it is 10,000 
 times as great as in pure water. We shall now make use of the indicator 
 methyl orange to examine solutions in which the concentration of H + is between 
 10~ 3 N and that of pure water, 10~ 7 N. Experiment. By a 10 fold dilution of 
 your 10- 3 A 7 HC1 solution, prepare 50 or 100 cc. 10~ 4 N HC1; and from this 
 prepare a 10~ 5 N solution. Test 10 cc. portions of these three solutions and of 
 water with 1 drop of methyl orange. Repeat the experiment with nitric or sulfuric 
 
 [23] 
 
acid, starting with your 10~ 3 A 7 " solution. Note. Save one of the lO' 1 N solutions 
 for later use, Paragraph 8. 
 
 6. Summarize your results with the two indicators in a table which shows the 
 color obtained at various concentrations of H + . Show your table to the 
 instructor. 
 
 7 '. The weak acid which w r e shall now proceed to study is acetic acid, 
 HC 2 H 3 O 2 . It is a white, crystalline, rather volatile solid which melts near room 
 temperature and is very soluble in water. The formula of its sodium salt, sodium 
 acetate, is NaC 2 H 3 O 2 . We shall abbreviate these formulas to HAc and NaAc, 
 respectively. Questions. How many mols of acetic acid are contained in 1 liter 
 of 6 N acid? How many mols of sodium acetate could be prepared from 
 this quantity of acetic acid? 
 
 8. Experiment. From the laboratory 6 N HAc prepare solutions which are 
 approximately N, 0.5 N, 0.05 N and 0.01 N. Place a 10 cc. portion of each of 
 the acetic acid solutions (N to .01 N) in a labelled test-tube, place 10 cc. distilled 
 water in another test-tube, and test each solution with methyl violet as in Para- 
 graph 3. For comparison, repeat the test with HC1 solutions of suitable 
 concentrations. (Note. In color comparisons it is not safe to trust the memory, 
 or even written descriptions.) Determine the lowest concentration of acetic acid 
 at which the color with methyl violet is distinctly different from that with water, 
 and estimate approximately the concentration of hydrogen ion in two of the acetic 
 acid solutions, say in the N and O.I N solutions. Repeat the experiment, but 
 use methyl orange instead of methyl violet, and estimate the concentration of H + 
 ion in the O.I N solution and in a more dilute solution. 
 
 9. The acetic acid in the solution must be present either in the form of ions, 
 H + and Ac~, or in the un-ionized form, HAc. From your estimate of the concen- 
 tration of H* in the 0.1 normal solution, calculate the fraction of the acetic acid 
 which is ionized, and the fraction which is un-ionized. The fraction of the acid 
 which is in the form of ions is called the degree of ionization. State also the 
 concentrations of acetate ion, Ac~, and of the un-ionized acid, HAC, in the 0.1 
 normal acetic acid solution. Is acetic acid a weak or a strong acid? 
 
 10. Calculations. The concentration of the ions in acetic acid solutions have 
 been determined by other methods more accurately than is possible by these color 
 experiments. The concentrations of hydrogen ion in these solutions at room 
 temperature are given in the following table: 
 
 Concentration Concentration Concentration Concentration of Degree of 
 acid. of H + . of Ac~. un-ionized HAc. Ionization. 
 
 1 N .004 N 
 
 0.5 AT .003 N 
 
 r\N .0013 N 
 
 O.OIN .0004 AT 
 
 Fill in the remaining columns of the table and show your table to the instructor 
 at once. 
 
 11. Memorize the fact that acetic acid and water are weak electrolytes. 
 Questions. Refer to Assignment 21, Paragraph 5, and state what you would 
 expect to happen if a solution containing Ac~ at high concentration were mixed 
 with a solution containing H + at high concentration. What solutions would you 
 mix to try this experiment, and how could you prove with an indicator that a 
 reaction had taken place? 
 
 12. The reaction between sodium acetate and hydrochloric acid solutions. 
 Experiment. Prepare some approximately half normal hydrochloric acid. 
 Measure out two 15 cc. portions in test-tubes, and add two or three drops of 
 methyl violet to each. Measure 10 cc. 1 N sodium acetate and add the solution, 
 
 [24] 
 
a few drops at a time, to one of the 0.5 N hydrochloric acid solutions. After 
 each addition shake the mixture, record the color and note the volume of the 
 sodium acetate solution added. Write the equation for the reaction that has 
 taken place between H + and Ac~. Questions. If solutions containing 0.10 mol 
 HC1 and 0.05 mol NaAc were mixed what substances would be present in the 
 resulting solution? What would be the concentration of each if the final volume 
 were (a) 1 liter, (b) 500 cc.? 
 
 13. The neutralization of NaOH solution by acetic acid. The concentration 
 of H + in an acetic acid solution is small in comparison with the total concentration 
 of acid in the solution, but the student must not jump to the conclusion that the 
 results of Assignments 5 and 6 on neutralization would have been materially 
 different if acetic acid had been used throughout instead of hydrochloric acid. 
 In the following experiment we shall study qualitatively the reaction between 
 NaOH solution and acetic acid; cf. the quantitative experiment in Assignment 5, 
 Paragraphs 4 and 5. Experiment. To 10 cc. 6 TV NaOH in a porcelain dish 
 add about 8 cc. 6 A r HAc. (Does the mixture become warm?) Add 1 drop of 
 phenolphthalein and continue to add acetic acid slowly until the solution becomes 
 colorless; test again with the indicator (since the phenolphthalein color fades 
 in a concentrated NaOH solution) and finally add about 2 cc. acid in excess. 
 Evaporate the solution until crystals of salt begin to separate. (Caution. Do 
 not evaporate the solution to dryness, since the NaAc may decompose.) Allow 
 the mixture to cool, collect some of the moist salt, NaAc 3 H 2 O, and dry it 
 between filter papers. Prove that the salt contains sodium by the flame test, and 
 acetate by warming with 6 TV sulfuric acid and noting the odor of acetic acid. 
 Question. What evidence is furnished by this experiment that the neutralization 
 reaction 
 
 HAc + NaOH == H 2 O + NaAc 
 
 takes place when solutions of HAc and NaOH are mixed? The concentration 
 of an acetic acid solution can be determined by titration with a known sodium 
 hydroxide solution when a suitable indicator is used, phenolphthalein in this case 
 (and if time permits the student may perform this titration). 
 
 14. We shall now examine the above reaction in the same way as we have 
 already done in Assignment 21, Paragraph 4, in the case of the reaction between 
 a strong acid and a strong base. Questions. Which of the four substances are 
 strong and which weak electrolytes? What substance is present at high concen- 
 tration in an acetic acid solution which is not present at high concentration in a 
 sodium acetate solution? What substance is present in a NaOH solution which 
 is not present at high concentration in sodium acetate solution? What substance 
 is present at high concentration in a sodium acetate solution which is not present 
 at high concentration in either of the initial solutions? Write the equation for 
 the main reaction, and show it to your instructor. The study of this reaction will 
 be continued in Assignment 25. V < 
 
 15. Problems, (i) 20.0 cc. acetic acid solution were found to neutralize the 
 same volume of NaOH solution as 16.0 cc. 0.50 TV H 2 SO 4 ; phenolphthalein was 
 the indicator in both titrations. What is the concentration of the acetic acid 
 solution (a) in mols per liter, (b) in equivalents per liter, and (c) in grams 
 per liter? 
 
 (2) Outline experiments to distinguish between 
 
 (a) 1.0 N HNO 3 and 0.10 N HNO 3 
 
 (b) A solution of nitric acid and one of acetic acid which give the same 
 bluish color with methyl violet. 
 
 [25] 
 
ASSIGNMENT 23 
 
 STRONG AND WEAK BASES. THE USE OF INDICATORS TO MEASURE 
 HYDROXIDE ION CONCENTRATION 
 
 1. In Assignment 23 we shall develop a method of measuring approximately, by 
 means of indicators, concentrations of hydroxide ion between normal and 10~ 7 
 normal, the concentration in pure water. This method will be used in studying 
 solutions of a typical weak base. Since there is throughout a close relation with 
 the preceding Assignment only brief direction will now be given. The student 
 is expected to make use of the discussion and Questions in Assignment 22 in 
 correlating the results of the two Assignments. 
 
 2. Experiment. Prepare solutions of sodium hydroxide which are approxi- 
 mately normal, 0.1 normal, 0.01 normal, and 0.001 normal. (State in your 
 note-book how you prepared these solutions.) To 10 cc. of each solution in a 
 test-tube, add 1 drop of a solution of the indicator, trinitrobenzol. Record the 
 color obtained in each case, and observe especially the most dilute solution that 
 gives a color with the indicator. Make a second series of observations using 1 a 
 larger amount of the indicator, say 6 drops, in each case. Note that by usirig 
 the different amounts of indicator you can determine approximately concentra- 
 tions of OH~ in one case between N and .01 N, and in the other between O.I N 
 and 0.001 N. 
 
 3. Repeat the experiment with potassium hydroxide solution, and compare the 
 colors obtained at each concentration. If sodium hydroxide in solution is 
 completely ionized, what conclusion can you draw with respect to potassium 
 hydroxide? What concentrations of hydroxide ion can be measured by means 
 of this indicator, trinitrobenzol? 
 
 4. Experiment. Prepare solutions of NaOH or KOH which are approximately 
 10- 4 N and 1O 5 N. Test 10 cc. portions of the 10~ 3 N, 10~ 4 N and 1Q- 5 N 
 solutions, and water with the indicator phenolphthalein. Repeat, using litmus 
 instead of phenolphthalein. Note. Since the quantity of alkali in a given volume 
 of these dilute solutions is extremely small it is evident that large errors in 
 concentration may result if the test-tubes and flasks are not thoroughly washed 
 with .distilled water before use. Check your results by preparing fresh portions 
 of the 10~ 3 , 10~ 4 and 10~ 5 N solutions and repeating the experiment. 
 
 5. Summarize your results in a table which shows the color obtained with 
 each indicator at various concentrations of hydroxide ion. Compare this table, 
 and the corresponding tph^e in Assignment 22, Paragraph 6, with the table given 
 by Hildebrand on page 181. 
 
 6. In the next experiment we shall determine the concentration of hydroxide 
 ion in solutions of the weak base, ammonium hydroxide, NH^OH. The concen- 
 trated laboratory reagent, cf. Assignment 7, Problem (j), is prepared by dissolving 
 the gas ammonia, NH 3 , in water until the solution is nearly saturated with NH 3 
 at room temperature. Write the equation for the formation of ammonium 
 hydroxide from ammonia and water. Ammonium sulfate, (NH 4 ).,SO4, is an 
 example of a salt of this base. Question. What is the concentration in mols per 
 liter and in equivalents per liter of ammonium ion, NH 4 + , and of sulfate ion in a 
 0.1 molal solution of ammonium sulfate? 
 
 7. Experiment. From the laboratory 6 TV NH 4 OH solution prepare solutions 
 which are approximately N, 0.1 N, 0.01 N and 0.001 N. By experiments with the 
 indicators trinitrobenzol and phenolphthalein (write out the details of these 
 experiments in your note book) estimate the concentration of OH~ in at least two 
 of these solutions. In each case give also the concentration of NH 4 + and of 
 un-ionized NH 4 OH, and calculate the degree of ionization. Note. Accurate 
 
 [261 
 
determinations show that at each concentration the degree of ionization of 
 ammonium hydroxide is almost the same as that of acetic acid at the same 
 concentration; see Assignment 22, Paragraph 10. 
 
 8. The reaction between ammonium chloride and sodium hydroxide solutions. 
 Predict what will happen when a solution of ammonium chloride is added to a 
 solution of sodium hydroxide; cf. Assignment 22, Paragraphs 11 and 12. Plan 
 an experiment to demonstrate this result and try the experiment. Write the 
 ionic equation for the reaction. 
 
 9. The neutralization of H^SO^ solution by ammonium hydroxide . Experiment. 
 To 10 cc. 6 N H 2 SO 4 in a" flask add 6 N NH 4 OH from a graduate until the 
 solution, after shaking, has a distinct odor of NH 3 . Evaporate in a porcelain 
 dish until a considerable quantity of salt has separated, cool the mixture, collect 
 some of the salt and dry it between filter papers. Prove that the salt contains 
 (a) sulfate, by dissolving some of it in water and adding barium chloride (to 
 precipitate BaSO 4 ), and (b) the ammonium radical, by warming with sodium 
 hydroxide solution and noting the odor of NH 3 . Give the experimental evidence 
 and reasoning in favor of the conclusion that, when solutions of a strong acid 
 and ammonium hydroxide solutions are mixed, the main reaction is: 
 
 H + + NH 4 OH = H 2 O + NH 4 * 
 The study of this reaction will be continued in Assignment 25. 
 
 10. Problems. (/) State how you would determine whether an unknown 
 solution is more acidic or more basic than water. (2) Suggest experiments to 
 distinguish between (a) 0.1 N KOH and 0.01 N KOH. (b) A solution of a 
 strong base and one of a weak base which have the same hydroxide ion 
 concentration. 
 
 ASSIGNMENT 24 
 RAPID REVERSIBLE REACTIONS AND EQUILIBRIUM 
 
 References. Hildebrand, Chapter XII, pages 155-171; Chapter XI, pages 145 
 
 and 148. 
 
 1. While certain reactions proceed to completion, as the transformation of 
 metallic copper into cuprous sulfide studied in Assignment 1, many reactions do 
 not. For example, when solutions containing equivalent amounts of a strong acid 
 and sodium acetate are mixed, cf. Assignment 22, Paragraph 12, the resulting 
 solution still contains about 1% of the reacting substances H + and Ac~, i. e., the 
 reaction H + -f- Ac" = HAc, although it takes place very rapidly, stops when about 
 99% of the possible amount of HAc has been formed. In the final solution, the 
 concentration of the three substances involved in the reaction are the same as in 
 the corresponding acetic acid solution. Similarly, if we had started with pure 
 acetic acid, which is un-ionized in the solid, liquid or gaseous state, and dissolved 
 it in water, the reaction HAc = H + + Ac~ would have taken place rapidly, but 
 only until about 1% of the acetic acid had been ionized. Question. From your 
 results in Assignment 22, what would be the concentrations of H + , Ac~ and HAc 
 in a solution made (a) by dissolving 1 mol of HAc to give a liter of solution, and 
 (b) by dissolving 1 mol of HC1 and 1 mol of NaAc to give a liter of solution? 
 
 2. The reaction just considered is an example of a rapid, reversible reaction, 
 and the three substances involved in this reaction are in equilibrium with each 
 other in the final solution. In general whenever it has been shown experimentally 
 that a reaction can be made to take place in both directions, i. e., is reversible, 
 then it may be concluded that under suitable experimental conditions a state of 
 equilibrium can be realized in which all the substances involved in the reaction 
 
 [27] 
 
are present together. For each set of experimental conditions there is a definite 
 state of equilibrium ; and when the experimental conditions are altered, e. g., by 
 changing the concentration of one or more of the substances involved or by 
 changing the temperature the reaction takes place in one direction or the other 
 until equilibrium is again established. The problem is to learn to predict what 
 will happen in any given case when the experimental conditions are altered. 
 
 3. The effect of changing the concentration of one of the substances involved 
 in an equilibrium. Experiment. Place two 10 cc. portions of N acetic acid in 
 two test-tubes, a 10 cc. portion of 0.1 N acetic acid in a third test-tube, and a 
 10 cc. portion of water in a fourth test-tube. To each solution add the same 
 number of drops of methyl violet solution. To one of the normal solutions add a 
 small amount of solid sodium acetate (or of 4 AT" solution) ; compare the colors 
 of the four solutions. Repeat the experiment with more dilute solutions of 
 acetic acid, using methyl orange instead of methyl violet. Question. What 
 conclusion can you draw with regard to the change in the concentration of 
 hydrogen ion in this experiment? What reaction must have taken place? When 
 equilibrium has again been established after the addition of the sodium acetate, 
 is the concentration of each of the substances H + , Ac" and HAc greater or less 
 than its concentration in the original acetic acid solution? State briefly how this 
 experiment illustrates the general statement : the effect of changing the concen- 
 tration of one of the substances involved in an equilibrium is to cause that reaction 
 to take place which tends to neutralize the change. The effect of increasing the 
 concentration of acetate ion could also have been predicted from the quantitative 
 statement of the Mars Law : (Concentration of H + ) (Concentration of Ac~) /^Con- 
 centration of HAc) = constant, when equilibrium has been established at a 
 definite temperature. 
 
 4. Outline an experiment to demonstrate that the reaction NH 4 + -|- OH~ = 
 NH 4 OH takes place when a solid ammonium salt is added to a solution of 
 ammonium hydroxide ; cf . your experiments in Assignment 23. Perform this 
 experiment, and explain how it illustrates the italicized statement in the preceding 
 Paragraph. 
 
 5. As another example of the effect of change of concentration upon equilibrium 
 we shall study the equilibrium between solid silver acetate and its ions. The 
 reversible reaction is: 
 
 AgAc ( solid )= Ag+ + Ac~ 
 
 Note. The solubility of silver acetate is 0.06 mol per liter at room temperature; 
 while it is much more soluble than a sparingly soluble salt, as silver chloride, 
 it is much less soluble than such salts as silver nitrate, sodium nitrate or sodium 
 acetate. Experiment. Prepare some solid silver acetate by adding to 25 cc. 
 4 N NaAc solution (free from chloride) * about 20 cc. 0.1 N AgNCX solution, 
 and shaking the mixture several times. Collect the solid on a filter paper, and dry 
 it by pressing between dry filter papers. Prepare a saturated solution by shaking 
 the solid with 10 cc. water at intervals for about 10 minutes. (The saturated 
 solution can be prepared somewhat more conveniently by warming the mixture to 
 40 or 50, but it must be cooled to room temperature before continuing the 
 experiment). Allow the solid to settle. Pour half of the clear saturated 
 solution into another test-tube and set it aside for later use, Paragraph 7. To the 
 remaining mixture of solid and saturated solution add 6 N HNO 3 drop by drop, 
 shaking the mixture after each drop is added. When the silver acetate has 
 dissolved, heat the solution nearly to boiling and note the odor. 
 
 Note. Place all silver residues, including any solution which contain an appre- 
 ciable amount of silver, in the jar maked "silver waste." 
 
 * If chloride is present it may be removed from NaAc solution by adding first some 
 AgNOs solution, shaking the mixture and filtering off the precipitated AgCl. The resulting 
 solution contains some NaNOs and should be used only in this Assignment. 
 
 [28] 
 
6. The disappearance of solid silver acetate on the addition of a strong acid 
 may be considered to be due to a shifting of the equilibrium 
 
 AgAc (solid) = Ag + + A<r 
 
 as a result of the establishment of the second equilibrium 
 
 H + + Ac- = HAc 
 
 When a drop of nitric acid is added the H + uses up Ac~ to form HAc; and 
 in order to again build up the concentration of Ac', and thus restore equilibrium, 
 some solid AgAc goes into solution. When more HNO 3 is added the same 
 processes are repeated, until finally all the silver acetate is dissolved. From this 
 discussion of the mechanism of the reaction it is evident that the result of the 
 experiment could have been predicted from the fact that acetic acid is a weak 
 acid. The principal substances that have disappeared are solid AgAc and H + , 
 and those that have formed are Ag + and HAc; and we may, therefore, write 
 for the main reaction the equation : 
 
 AgAc (solid) + H + == HAc + Ac-. 
 
 The same equation is obtained by adding together the two equations considered 
 ' in discussing the mechanism of the reaction. It is important, however, to realize 
 that in such a case as this no single equation can represent all that has happened 
 in the reaction. Thus, the concentration of Ac~, although small throughout the 
 experiment, is much smaller in the final solution than in the initial AgAc solution ; 
 this small decrease is due to the reaction H+ + Ac~ = HAc, but is not taken into 
 consideration in the equation which we have written for the main reaction. 
 
 7. Predict what will happen when solid sodium acetate is added to a saturated 
 solution of silver acetate. Experiment. Test your answer by adding a small 
 amount of 4 AT NaAc to the clear saturated solution of silver acetate which you 
 have just prepared. Question. How does this experiment illustrate the italicized 
 statement in Paragraph 3? Re-word that statement to make it apply specifically 
 to the change of solubility of a salt when the concentration of one of the ions of the 
 salt is changed. What will happen when solid silver nitrate is added to a saturated 
 solution of silver acetate? 
 
 8. Problems. (/) For each of the following reactions: 
 
 (a) H 2 O (liquid) =H 2 O (gas) 
 
 (b) H 2 O (solid) =H 2 O (liquid) 
 
 (c) NH 3 (in solution) = NH 3 (gas) 
 
 (d) Ag + + Cl- AgCl (solid). (Note the solubility of AgCl is 1.0 X 10' 5 
 mols per liter) ; 
 
 outline an experiment to show that the reaction can take place as written and a 
 second experiment to show that the reverse reaction can also take place. 
 
 (2) In each case in Problem (/) point out the conditions under which equi- 
 librium can be realized. A satisfactory answer in (a) is the following: whenever 
 water and water vapor are present together in a closed vessel equilibrium is 
 quickly established, and remains unchanged as long as the temperature remains 
 constant; when the closed vessel contains only water and water vapor the 
 pressure at equilibrium is 1 atmosphere when the temperature is 100 ; cf . 
 Assignment 2, Paragraph 10, for the pressure of water vapor at room temperature. 
 
 (j) Give an example of a reaction which can be shown to be reversible at high 
 temperature, but which proceeds so slowly at room temperature that equilibrium 
 will not be realized in several months. Note that all the substances involved in 
 this reaction can exist together at room temperature but that such a fact does not 
 constitute a proof that equilibrium has been established. 
 
 (4) From your table in Assignment 22, Paragraph 10, state what fraction of 
 
 [29] 
 
the acetic acid is ionized in the 0.1 N and 0.01 N solutions. What reaction takes 
 place when 1 liter of 0.1 N solution is diluted to 10 liters of 0.01 N solution? 
 
 (5) State how the experimental result in Problem (4) could have been pre- 
 dicted from kinetic considerations, on the assumption that the two opposing 
 reactions, HAc = H + -j- Ac~ and H + -f- Ac~ = HAc, are actually taking place at 
 equilibrium. 
 
 (6) When one of the substances involved in an equilibrium is a solid, what 
 is the effect of adding more of this solid to an equilibrium mixture? Does the 
 solubility of a salt at room temperature depend on the amount of the salt in 
 contact with the saturated solution? Give another example of a reversible 
 reaction involving a solid substance. 
 
 ASSIGNMENT 25 
 THE REVERSIBILITY OF NEUTRALIZATION REACTIONS. HYDROLYSIS 
 
 References. Hildebrand, Chapter XIII, pages 185-192; as a review, read pages 
 
 173 and 180-183. 
 
 1. In earlier Assignments we have studied various examples of neutralization 
 reactions and have shown that the main reactions that take place are represented 
 by the following equations : 
 
 (a) H + + OH" = H 2 O when a strong acid reacts with a strong base. 
 
 (b) HA + OH~ = H 2 O + A~ when a weak acid reacts with a strong base. 
 
 (c) H + + BOH = H 2 O + B + when a strong acid reacts with a weak base. 
 
 In the present Assignment we shall find that these reactions are reversible and 
 we shall study the equilibrium that is established in each case. When the 
 Assignment is finished the student should be able, from a knowledge of the 
 relative strengths of various acids, and of various bases, to predict the approxi- 
 mate concentrations of the various substances at equilibrium in any given 
 neutralization reaction. 
 
 2. We wish to examine experimentally the properties of the solution that is 
 obtained when equivalent amounts of a given acid and base react. To do this we 
 shall take advantage of the fact that a solution of a pure salt is identical with the 
 solution prepared by mixing exactly equivalent amounts of the corresponding 
 acid and base. In other words we shall make use of the reaction which is the 
 reverse of neutralization in establishing the equilibrium state that we desire to 
 study. The experimental work will consist in determining by means of indicators 
 the approximate concentration of H + or OH~ in solutions of various salts. 
 Note. In the following experiments it is important that all test-tubes should be 
 clean; rinse each several times with distilled water before using it. 
 
 3. Experiment. Test 10 cc. portions of N NaCl solution and distilled water by 
 adding to each the same number of drops of litmus solution. Repeat using 
 phenolphthalein and methyl orange. Question. What conclusions can you draw 
 with regard to the relative concentration of H + or OH~ in sodium chloride solution 
 and in water? Write the ionic equation for the neutralization of NaOH by HC1 
 solutions, and state whether the Na + and Cl" are involved in the reaction. In 
 pure water concentrations of H + and OH~ are the same and equal to 10~ 7 
 mols per liter, what are the concentrations of Na + , Cl~, H + and OH~ in the 
 solution made (a) by dissolving 1 mol NaCl in water to give a liter of solution, 
 and (b) by dissolving exactly 1 mol of HC1 and 1 mol of NaOH to give a liter 
 of solution? 
 
 4. Experiment. Test 10 cc. portions of 4 N NaAc and distilled water first 
 
 [301 
 
with litmus solution and then with phenolphthalein. Estimate approximately the 
 concentration of OH" in the 4 N NaAc solution. Questions. Is the concentration 
 of OH" in 4 N NaAc solution greater or less than in pure water? From the fact 
 that the solution must be identical with that obtained by mixing exactly equivalent 
 amounts of acetic acid and sodium hydroxide, what conclusions can you draw 
 with regard to the completeness of the neutralization reaction between HAc and 
 OH"? To account for the formation of OH" when NaAc is dissolved in water, 
 what reaction must take place? What are the approximate concentrations of 
 Na + , Ac", HAc and OH- in the 4 N NaAc solution? 
 
 5. It will be observed that the main reaction which accounts for the neutrali- 
 zation of HAc by NaOH, 
 
 HAc + OH- + (Na + ) = Ac- + H 2 O + (Na + ) 
 
 or the reverse of this reaction considered in the preceding paragraph, involves 
 the disappearance of one weak electrolyte and the formation of another. Both 
 the OH" and the Ac" are competing for the H + . From the fact that water is a 
 far weaker electrolyte than acetic acid, it could have been predicted that when 
 equilibrium is reached the OH" will have taken nearly all the H + away from 
 the Ac" ; i. e., the concentration of the HAc and OH" will be very small and that 
 of the Ac" large. Question. From your table of the concentrations of H + in 
 acetic acid solutions, Assignment 22, what is the ratio of the concentration of H + 
 in N acetic acid to that in pure water? 
 
 6. As indicated in the preceding paragraph, the mechanism of this reaction 
 may be considered to be parallel to that of the dissolving of silver acetate in a 
 solution of a strong acid, Assignment 24. When an acetic acid solution is 
 neutralized by NaOH solution, the OH" unites with H + which is present at small 
 concentration in the acetic acid solution : 
 
 H + + OH- = H 2 O 
 
 and, in order to again restore equilibrium, the HAc dissociates, 
 
 HAc = H + + Ac- 
 
 These two reactions continue to take place simultaneously until equilibrium is 
 established. The sum of the two gives the equation which we wrote for the main 
 reaction. 
 
 7. Discuss in a similar way the mechanism of the reaction between Ac" and 
 H 2 O which takes place when sodium acetate is dissolved in water. This, of 
 course, involves the reverse of the two reactions written in Paragraph 6. Write 
 the equation for the main reaction in this case. 
 
 8. The reaction which is the reverse of neutralization is called hydrolysis. 
 In the reaction which we have just considered a small amount of the acetate ion 
 is said to be hydrolyzed. Question. Why would you expect the Ac" to be 
 hydrolyzed slightly while the Cl~ is not hydrolyzed at all? 
 
 9. The dissociation of HCXV, bicarbonate ion, into H + and carbonate ion, 
 CO 3 ", is far less than the dissociation of acetic acid. Questions. What reaction 
 would you expect to take place between CO 3 ~~ and H 2 O when Na 2 CO 3 is 
 dissolved in water? Which would you expect to be the more basic, a solution of 
 Na 2 CO 3 or a solution of NaAc of the same molal concentration? Test your 
 answer by the following experiment 
 
 10. Experiment. Estimate the approximate concentration of OH" in a 0.5 M 
 Na 2 CO 3 solution by testing 10 cc. portions with litmus, phenolphthalein and 
 trinitrobenzol. The main reaction for the hydrolysis of CO 3 ~ " is 
 
 C0 3 -- + H 2 = HCO 8 - + OH- 
 
 and the concentration of OH~ in a 0.5 M Na 2 CO 3 solution is about 0.01 M. 
 Questions. What would be the concentration of OH- if the CO 3 - " were completely 
 
 [31] 
 
hydrolyzed? What fraction of the CO 3 ~~ has been hydrolyzed? What would be 
 the approximate concentrations in mols per liter of Na + , CO 3 ~~, HCO 3 ~ and OH" 
 in a solution prepared by dissolving 0.05 mol NaHCO 3 and 0.5 mol NaOH to 
 give 1 liter of solution? 
 
 11. Experiment. Test a 10 cc. portion of 4 AT NH 4 C1 with litmus solution. 
 Compare the color with that obtained by adding the same amount of litmus 
 solution to (a) 10 cc. water and (b) 10 cc. water containing 1 drop of 6 N HC1. 
 Questions. How do you account for the fact that a solution of NH 4 C1 is slightly 
 acid while that of NaAc is slightly basic. Write the main reaction for the 
 hydrolysis of the NH 4 + , and the two reactions which may be used to explain the 
 mechanism of the reaction. The concentration of H + in 4 N NH 4 C1 is about 
 4 X 10~ 5 N. What is the concentration of each of the principal substances 
 present in a solution made by dissolving 4 mols of HC1 and 4 mols of NH 4 OH 
 to form a liter of solution? 
 
 12. Problems. (1) What is the concentration of H + in N NaOH? 
 Note. Consult Hildebrand, page 180. 
 
 (2) From the results of the experiments in this Assignment and the table of 
 indicators (Hildebrand, page 181) select one or more indicators which might be 
 used to determine the end-point (i e., to determine when equivalent amounts of 
 acid and base are present) in each of the following titrations : 
 
 (a) NaOH solution with HC1 or HNO 3 solution. 
 
 (b) HAc solution with NaOH solution. 
 
 (c) NaHCO 3 solution with NaOH solution. 
 
 (d) NH 4 OH solution with HC1 solution. 
 Note. Compare Assignment 6, Paragraph 3. 
 
 (5) An ammonium acetate solution gives the same colors with indicators as 
 does water : What hydrolysis reactions must occur when solid NH 4 Ac is dissolved 
 in water? State what substances are present in the solution. 
 
 (4) Point out parallelisms between the strengths of acids and bases and the 
 hydrolysis of the corresponding ions. 
 
 32 
 
SECTION III 
 REACTIONS OF IONS 
 
 ASSIGNMENT 31 
 
 THE PROPERTIES OF SODIUM, POTASSIUM AND AMMONIUM IONS. 
 TESTS FOR CHLORIDE, SULFATE AND NITRATE IONS 
 
 Reference, for Assignments 31 and all later Assignments: A Standard Text on 
 
 Inorganic Chemistry. 
 
 1. In this Assignment we shall study the properties of the ions of the common 
 acids and bases. 
 
 2. Aside from the physical properties, such as color and taste, the properties 
 of an ion depend on its behavior toward other substances. The study of the 
 chemistry of an ion thus consists in determining whether or not various substances 
 react with the ion, and in studying the reactions that do take place. For each 
 reaction it is necessary to know : 
 
 (1) the products of the reaction, 
 
 (2) whether the reaction takes place rapidly or slowly, 
 
 (j) whether the reverse reaction can be made to take place, and 
 
 (4) the nature of the equilibrium if the reaction is rapid and reversible. 
 
 In this course the task is greatly simplified on account of the following consid- 
 erations : Nearly all the reactions studied take place very rapidly ; this is 
 characteristic of ionic reactions, such as ionization, precipitation, neutralization, 
 etc. Also the number of different types of reactions is not large ; and, when for 
 one type a single example is understood, additional examples should present no 
 real difficulty. Finally an equilibrium state is reached in many cases, and a 
 knowledge of the principles of equilibrium enables the student to predict what 
 will happen when there is a given change in the experimental conditions. While 
 the number of facts to be remembered is thus greatly reduced, it is necssary, 
 however, to memorize such facts as; the formulas of substances, the relative 
 solubility of salts, the relative volatility of various solid or liquid substances and 
 the relative strength of weak electrolytes. 
 
 3. Properties of sodium ion. The chemistry of Na + is extremely simple and 
 may be summarized by saying that, at ordinary temperatures, with a few 
 uncommon exceptions, all of its compounds are non-volatile, readily soluble and 
 strong electrolytes. 
 
 4. Flame test for sodium. Experiment. Clean an iron wire by dipping it into 
 dilute HC1 solution in a test-tube and holding the wire in the flame until it gives 
 only a faint yellow color. (Note. Do not contaminate your HC1 solution hy 
 dipping the iron wire into the bottle.) Try the flame test with a solid sodium salt, 
 the laboratory solution of a sodium salt, a very dilute solution and distilled 
 water. The test is extremely delicate and substances which do not contain 
 measurable amounts of sodium usually give a slight yellow color for a short 
 time. Therefore, before concluding that sodium is present in an "unknown" 
 it is well to make comparative tests with known solutions (/) free from sodium 
 and (2) the same solution containing a small amount of sodium salt. Such 
 experiments are called blank tests. 
 
 5. Properties of potassium ion. What is the relation of K and Na in the 
 Periodic System? Potassium ion resembles .sodium ion very closely. One 
 difference, however, is the precipitation of potassium cobaltinitrite when cobalti- 
 nitrite ion, Co(NO 2 ) fi is added to a solution containing K + . Experiments. To 
 
 [ 33 ] 
 
5 cc. of a dilute, neutral solution of a potassium salt add a few drops of acetic 
 acid and 2 to 5 cc. of the sodium cobaltinitrite reagent ; let the mixture stand 
 10 minutes. 
 
 6. Flame test for potassium. Try the flame test with solid KC1, KC1 solution, 
 KC1 solution containing a small amount of Nad, Nad solution, and NaCl solution 
 containing a small amount of KC1. Use the blue glasses to shut off the yellow 
 color of the sodium flame and practice until the presence or absence of both 
 potassium and sodium can be determined. Note. Since the flame test depends 
 upon the amount of solid present it is well to concentrate a dilute solution by 
 evaporation before making the test. Chlorides give more satisfactory flame tests 
 than sulfates or oxides, which is due to the relatively greater volatility of the 
 chlorides. 
 
 7. Properties of ammonium Ion. The ammonium ion resembles sodium ion and 
 potassium ion in that with few exceptions its compounds are soluble, strong 
 electrolytes. Like the potassium ion it forms a precipitate with Co(NO 2 ) 6 . 
 Repeat the experiment in Paragraph 4, using NH 4 + instead of K + . Ammonium 
 compounds give no flame test. The chemistry of NH 4 + is made somewhat more 
 complicated by the fact that NH 4 OH not only is a weak base, but is decomposed 
 easily into NH 3 and H 2 O. NH 3 , ammonia, enters into many reactions which 
 will be considered in later Assignments. Experiment. Note the odor of the 
 
 6 N NH 4 OH laboratory reagent; prepare some dilute solutions of NH 4 OH, e. g., 
 0.1 A 7 " etc., heat gently about 10 cc. of each solution (moving the test-tube back 
 and forth through the flame) and note the odor from time to time; decide what 
 is the most dilute solution you can detect by means of the odor of NH 3 gas. 
 Determine by experiment whether you can detect the odor of NH 3 gas when a 
 solution of NH 4 C1 or (NH 4 ) 2 SO 4 is boiled. Suggest a method of testing for 
 ammonium ion based on the formation first of NH 4 OH in the solution and then 
 of NH 3 gas. Write equations. Experiment. Try your method with a dilute 
 solution of an ammonium salt. 
 
 8. Volatility of ammonium compounds. Ammonium compounds are much 
 more volatile than those of sodium and potassium. Experiment. Evaporate a 
 small amount of (a) NH 4 C1 solution and (b) a mixture of NH 4 C1 and KC1 
 solutions to dryness in a porcelain dish and heat until fumes are no longer given 
 off. Suggest a method other than the flame test of determining the presence of 
 K + in a solution which contains both K + and -NH 4 + . 
 
 9. Properties of H + and OH~. Aside from the tendency of H + and OH" to 
 unite with each other to form water each of these ions reacts with many other 
 ions to form weak electrolytes. We have already noted the reaction between H + 
 and Ac~ and OH~ and NH 4 + . Acids are more volatile than their salts : the relative 
 volatility of the common acids will be considered in Assignment 34. The OH~ 
 also forms precipitates with many metallic ions. Question. How would you 
 test for (and estimate the concentration of) H + or OH" in an unknown solution? 
 
 10. Although the ions NO 8 ~, Cl~ and SO 4 " will not be considered in detail in 
 this Assignment, certain of their reactions will be studied in order that their 
 presence in solutions may be detected. All nitrates are soluble, but a few 
 chlorides and sulfates are sparingly soluble. 
 
 11. Test for chloride ion. Experiment. To 1 cc. portions of solutions of NaCl, 
 Ba(OH) 2 , Na 2 CO 3 and Na 2 SO 4 add 10 cc. water, shake and add a few drops 
 of AgNO 3 solution. In each case, record your observations, noting the color of 
 the precipitates. Now add HNO 3 until each solution is acid. From your results 
 state how you would test an unknown solution for chloride ion. Note. Ba(OH) 2 
 was used instead of NaOH because Cl~ is usually present in the laboratory NaOH 
 solutions. Experiment. Test a dilute solution of your NaOH for Cl~. 
 
 [34] 
 
12. Test for Sulfate ion. Experiment. To dilute solutions of NaCl, NaOH, 
 NaXO 3 and Na 2 SO 4 add a few drops of Ba(NO 3 ) 2 solution. Now add HNO 3 
 until each solution is distinctly acid. From your results state how you would 
 test an unknown solution for SO 4 ~~. 
 
 13. Test for nitrate ion. Try the following experiment with a solution which 
 contains nitrate ion and with one free from nitrate ion, first acidifying with 
 6 N H 2 SO 4 if the solution is alkaline. To about 2 cc. of the solution to be 
 tested for nitrate ion add in excess a solution of ferrous sulfate, FeSO 4 (or of 
 ferrous ammonium sulfate, a double salt of ferrous sulfate and ammonium 
 sulfate), filter if there is a precipitate, hold the test tube in a slanting position, 
 and pour carefully down the side of the test-tube (from a small beaker) 2 or 
 3 cc. concentrated sulfuric acid. The concentrated acid sinks to the bottom 
 of the tube and a dark brown ring forms on its surface when nitrate is present. 
 Note. The brown substance decomposes rapidly when the mixture is hot. 
 If the result of the test is negative and the test-tube feels hot to the hand (owing 
 to the heat liberated when the concentrated H 2 SO 4 mixes with the solution), 
 repeat the test more carefully. 
 
 14. Make up solutions of various concentrations of Cl~ and try experiments 
 to determine if your test can be used to decide whether the concentration of 
 Cl~ is large or small. Repeat for the nitrate and sulfate tests. 
 
 15. Analyses Nos. A and B. Analyze the unknowns for the four positive and 
 four negative constituents considered in this Assignment. Try to distinguish 
 between large amounts, small amounts, and traces. 
 
 16. Problems. (/) What is the valence of each of the 4 positive and 4 negative 
 ions considered in this Assignment? Write the formulas of the sixteen possible 
 compounds, each of which contains one positive and one negative constituent 
 (cf. Hildebrand, pages 90-91). Name each of the sixteen substances, state 
 whether it is a solid, liquid or gas at room temperature, whether it is readily 
 soluble in water or not, and whether it is a strong or weak electrolyte. 
 
 (2) What is the valence of sodium in (a) solid sodium sulfate, (b) metallic 
 sodium? 
 
 (j) A solid unknown was dissolved in wa f er and the solution was found to give 
 distinct tests for Na + , K + , Cl~ and SO 4 ~ ~. What conclusions can you draw with 
 regard to the nature of the solid salts in the original mixture? 
 
 (4) Outline experiments to decide whether or not each of the following 
 contains the impurity named : 
 
 (a) Barium chloride impurity in the barium hydroxide solution. 
 
 (b) Potassium chloride in solid sodium chloride. 
 
 (c) Potassium chloride in solid ammonium chloride. 
 
 (d) Ammonium chloride in solid potassium chloride. 
 
 (e) Sodium nitrate in solid sodium sulfate. 
 
 (5) In each of the following it is assumed that two solutions are mixed each 
 of which contains one of the substances at moderate concentration, say 0.1 M . 
 Mark the cases in which no reaction takes place. In the others write the equation 
 for the reaction. In (a), (b), (c) state what solutions you would use in the 
 experiment. 
 
 (a) Ac-andCl- (d) NH 4 OH and K* (g) K + and Ac~ 
 
 (b) NH t + and Na + (e) H + and Ac- (h) Ag + and Cl- 
 
 (c) NH 4 OHandH + (/) K + and OH- (i) Ba ++ and Cl~ 
 
 F35 1 
 
ASSIGNMENT 32 
 CALCIUM ION 
 
 References. Hildebrand, pages 175-179, 192-193 and 207. 
 
 1. In Assignment 32 we shall study the chemistry of metallic calcium and of 
 calcium ion, including its formation from metallic calcium. We shall find that 
 certain compounds of this element differ from the corresponding compounds of 
 sodium, potassium, and ammonium in that they dissolve to a much smaller extent 
 in water; and we shall study the equilibrium between these solids and their 
 saturated solutions. 
 
 2. Note. In future write equations for all reactions. 
 
 3. Obtain from the, office a piece of metallic calcium. Describe its properties as 
 far as you can observe them by physical examination. In what respects does 
 calcium show the physical properties of a metal? 
 
 4. Reaction between calcium and water. Experiment. Drop the calcium 
 into about 20 cc. of water. Stir the mixture and warm it gently until the metal 
 has disappeared. Test the solution for OH-. What reaction has occurred ? The 
 white solid formed, calcium hydroxide, Ca(OH) 2 , is a strong base but only 
 moderately soluble in water. Cork the test-tube and save the mixture for use 
 in a later experiment. Question. Name two other metals that react readily with 
 water. What reactions occur? 
 
 5. Calcium hydroxide. If solid calcium hydroxide is shaken with water until 
 a saturated solution is obtained an equilibrium between solid Ca(OH) 2 and the 
 ions Ca + * and OH~ is established. The equation for the reaction that has taken 
 place is 
 
 Ca (OH ) 2 (solid) = Ca ++ + 2OH- 
 
 The solubility of Ca(OH) 2 at room temperature is 0.02 mols per liter. Questions 
 What is the concentration of calcium ion, of hydroxide ion, (/) in mols per 
 liter, and (2) in equivalents per liter? Is the reaction just considered reversible? 
 What solutions would you mix in order to prepare solid Ca(OH) 2 ? Try the 
 experiment. 
 
 6. Which contains the larger concentration of OH~, a saturated solution of 
 Ca(OH) 2 or 1 AT NH 4 OH (see Assignment 23)? State what you think will 
 happen when a dilute solution of calcium chloride is made alkaline with ammonium 
 hydroxide. Try the experiment. Explain the result if your prediction was 
 incorrect. 
 
 7. State what would happen if a solution of calcium hydroxide were treated 
 with (i) hydrochloric acid, (2) ammonium chloride. Predict what would 
 happen if solid Ca(OH) 2 were also present in each case. Experiment. Test 
 your answers by experiments with portions of the mixture prepared in Par- 
 agraph 4. Questions. How would you use calcium hydroxide to test a solution for 
 ammonium ion? 
 
 8. Calcium carbonate. The solubility of CaCO 3 in water Lt room temperature 
 is very small, viz., 0.00013 mols per liter. Question. What is the concentration 
 of calcium ion, of carbonate ion, in this saturated solution, (i) in mols per 
 liter, (2) in equivalents per liter? What solutions would you mix in order to 
 form a precipitate of calcium carbonate? Try the experiment. Write the 
 simplest ionic equation for the reaction. Continue the experiment with very 
 dilute solutions of a calcium salt in order to determine if this reaction can be 
 used as a delicate test for calcium ion. In the more dilute solutions, if a pre- 
 cipitate does not form at once, heat to boiling and let stand 10 minutes. If the 
 liquid is turbid compared with water a precipitate has formed. 
 
 [361 
 
9. Experiment. Dilute 5 cc. N CaCL with about 100 cc. water, add about 8cc. 
 N Na 2 CO 3 solution, heat the mixture to boiling, filter, wash the precipitate and 
 reject the wash water. Consider what substances may be present in the filtrate. 
 Test for chloride ion and for calcium ion. Which reagent was present in excess, 
 calcium chloride, or sodium carbonate? 
 
 10. Experiment. Treat a portion of the CaCO 3 precipitate with hydrochloric 
 acid solution. Repeat the experiment with nitric acid. The gas given 
 off is carbon dioxide, CO 2 . What substances are present in the final solution in 
 each case after the solution has been boiled to expel the CO 2 ? Note. CO 2 is 
 moderately soluble in cold water and the solution contains carbonic acid, 
 H 2 CO 3 , which is a weak acid. CO 2 gas will not be evolved until the concentra- 
 tion of carbonic acid in the solution becomes sufficiently high. Questions. How 
 is the equilibrium between solid CaCO 3 and its ions disturbed by the addition 
 of a strong acid? What is the analogy between this example and the action 
 of an acid on solid Ca(OH) 2 , and on solid silver acetate? What is the reason 
 for the difference between these results . and that obtained when dilute nitric 
 acid is added to silver chloride or barium sulfate? 
 
 11. Test for carbonate. Experiment. Fit a test-tube with stopper and delivery 
 tube. Place a small amount of any carbonate in the tube and add a few cc. of 
 6 N H 2 SO 4 . Insert the stopper and pass the gas given off into a clear solution of 
 Ca(OH) 2 .* If only a small amount of carbonate is present the CXX will not pass 
 over into the Ca(OH) 2 solution. A satisfactory test may be obtained in this case 
 by placing the test-tube in a beaker of boiling water, or by adding a small piece of 
 zinc to the cold mixture in the test-tube. The CO 2 is carried over with the 
 dissolved air and steam in the first method, and with the hydrogen gas in the 
 second. 
 
 12. Calcium oxalate. Experiment. Prepare some CaC 2 O 4 by treating a 
 solution containing calcium ion with ammonium oxalate solution. Test the 
 action of excess of a strong acid, as HC1 or HNO 3 on CaC 2 O 4 . Question. Is 
 oxalic acid, H 2 C 2 O 4 , a weak or a strong acid? 
 
 13. Experiment. Determine if the precipitation of CaC 2 O 4 is a delicate 
 test for calcium by treating very dilute solutions of Ca ++ with ammonium oxalate 
 solution and a little ammonium hydroxide (to make sure that the mixture is not 
 acid). Observe the same precautions as in the test for calcium by means of 
 CO 3 . Note. The solubility of calcium oxalate in water is even less than that 
 of calcium carbonate. 
 
 14. Calcium sulfate. Experiment. To 10 cc. normal CaCl 2 solution add 
 2 cc. 6 normal H 2 SO 4 . If no precipitate appears at once, heat the solution 
 gently, and let it stand. Filter. Test a portion of the filtrate for Ca + + by adding 
 NH 4 OH until the solution is no longer acid, and then (NH 4 ) 2 CO 3 solution, and 
 warming the mixture. Test another portion for SO 4 ~~. What conclusion do you 
 draw with regard to the solubility of CaSO 4 in water? Which is the less 
 soluble, (i) CaSO 4 or CaCO 3 , (2) CaSO 4 or BaSO 4 ? 
 
 15. Predict what will take place when solid CaSO 4 is heated with excess 
 Na 2 CO 3 solution? Experiment. Test your prediction by boiling some solid 
 CaSO 4 with normal Na 2 CO s solution. Test the filtrate for SO 4 ~. Wash the 
 precipitate with water, and test it for carbonate. 
 
 16. Experiment. Test whether the reverse reaction will take place, by heating 
 solid CaCO 3 with sodium sulfate solution, filtering and testing the precipitate and 
 filtrate. 
 
 ^Instead of Ca(OH) 2 solution, Ba(OH)s. solution may be used; BaCO 3 , like CaCOs is 
 a difficultly soluble substance. 
 
 [37] 
 
17. Flame test for calcium. Experiment. Determine whether calcium salts 
 give a characteristic flame test. Can you distinguish the flame with a calcium 
 salt from that with sodium or potassium salts? 
 
 18. Note. The solubilities of many common salts at 18 are listed on the inside 
 cover page of Alexander Smith's text-book. Do not try to remember the 
 actual solubilities, but make lists of the readily soluble and difficultly soluble 
 salts of each metal studied and memorize these lists. Arrange the compounds ot 
 calcium according to their solubilities in water, distinguishing readily soluble, 
 moderately soluble, and difficultly soluble substances. Point out the compounds 
 which are much more soluble in dilute hydrochloric or nitric acids than in water. 
 
 19. Give as many methods as you can of testing for Ca ++ in a solution 
 which is known to contain H + , Na + , K + and NH 4 + . 
 
 20. Problems, (i) What is the minimum volume of water needed at 18 to 
 dissolve i gram of calcium carbonate? 
 
 (2) Can the following substances be present at moderate concentrations in 
 the same solution? If not, what is formed? 
 
 (a) H + andNO 3 - (e) H + and C 2 O 4 - - 
 
 (0) H + and OH- (/) Ca ++ and CO 3 - - 
 
 (c) H + and SO 4 - - (g) Ca ++ and NO 3 ~ 
 
 (rf) H + and CO 3 - - (h) Ca ++ and NH 4 OH 
 
 (j) How would you prepare: 
 
 (a) Solid CaCO 3 from solid Ca(OH) ? 
 
 (b) Solid CaSO 4 from solid CaCO 3 ?~ 
 
 (c) Solid CaC 2 O t from solid CaSO 4 ? 
 
 Write equations for all reactions, and point out what equilibria are involved. 
 
 (4) What is the position of calcium in the Periodic System? Name the 
 elements in the alkaline earth group. Write a brief note illustrating gradations 
 of properties in this group, after examining the table given by Hildebrand on 
 page 268. 
 
 ASSIGNMENT 33 
 CARBONATE ION, BICARBONATE ION, AND CARBONIC ACID 
 
 1. This assignment, in which we shall study the chemistry of carbonic acid and 
 its ions, is introduced, not only because this subject is an extremely important 
 one, but also because it serves to illustrate how a large number of facts can be 
 correlated when the equilibria involved in a few simple, reversible reactions are 
 understood. Carbonic acid, H 2 CO 3 , is a weak, dibasic acid; and, like other weak 
 acids which contain more than one acidic hydrogen in the molecule, it ionizes in 
 steps, to form (/) bicarbonate ion, HCO 3 ~, and (2) carbonate ion, CO 3 ~~, thus: 
 
 H.CO, == H + + HCCV (1) 
 
 HCO 3 - = H + + CO 3 -- (2) 
 
 In order to understand what happens in the various reactions considered, it is 
 necessary to know the relative concentrations, in the initial and final solutions, 
 of the substances involved in these two rapid reversible reactions. Our problem, 
 then, is to determine how the concentrations of H + (or OH"), H 2 COo, HCO 3 ~ 
 and CO 3 ~~ at equilibrium in a solution vary as the experimental conditions are 
 changed. 
 
 2. Obtain at the office a large and small two-holed rubber stopper, two pieces 
 of rubber tubing and a thistle tube for use in the experiment in paragraph 3 arid 
 a hard glass test-tube for use in the experiment in Paragraph 12. 
 
 [381 
 
3. Preparation and properties of carbon dioxide. Experiment. Make a carbon 
 dioxide generator by fitting your half-liter flask with the thistle tube 
 (passing through the stopper nearly to the bottom of the flask), and an outlet 
 tube bent at right angles. Make a wash bottle for washing the CO 2 gas by 
 equipping a small flask with an entry tube extending nearly to the bottom of the 
 flask, and a delivery tube. Fill the small flask about half full of water. Place 
 in the generator a few lumps (about 5 grams) of limestone, cover with water, 
 and add hydrochloric acid, a little at a time, through the thistle tube. 
 
 Note the color and the odor of the gas. Calculate from the molecular weights 
 the relative densities of carbon dioxide and oxygen, and of carbon dioxide and, 
 nitrogen, at the same temperature and pressure. Question. Is carbon dioxide 
 denser or lighter than air? Experiment. Collect some carbon dioxide in a 
 test-tube by displacement of air; and determine if it is inflammable, and if it 
 supports combustion. 
 
 4. Suggest an experiment to prove that carbon dioxide is moderately soluble in 
 water at room temperature. Try the experiment. When CO 2 dissolves in water 
 the reaction is 
 
 CO 2 (gas) + H 2 O = H 2 CO 3 (in solution). 
 
 Prove by an experiment that the reverse reaction can take place. Under what 
 conditions is there an equilibrium? Note. At room temperature a solution in 
 equilibrium with CO 2 gas at 1 atmosphere pressure contains about 0.04 mols 
 H 2 COo in 1 liter. Question. How is this equilibrium altered by an increase of 
 temperature ? 
 
 5. We shall now consider what the principal substances are in a saturated 
 carbonic acid solution. Experiment. Prepare a small quantity of nearly 
 saturated solution of CO 2 and test it with the indicators, phenolphthalein. litmus, 
 methyl orange, and methyl violet; for comparison repeat indicator tests with 
 water and with an acid solution of known H + concentration. What conclusion 
 can you draw with regard to the concentration of H + in the saturated CO 2 
 solution? Note. The concentration of H + in this solution has been determined 
 by other methods to be about 0.0001 N (W~ 4 N). From this value for the 
 concentration of H + , the concentration of OH" in the same solution can be 
 calculated to be 10 ~ 10 N; explain how this calculation is made, cf . Hildebrand, 
 page 180. 
 
 6. Experiment. Test the action of carbonic acid on Ca ++ by passing CO 2 
 gas into a dilute solution of CaCl 2 . What conclusion can you draw with regard 
 to the concentration of CO 3 ~~ in carbonic acid solution? The concentration of 
 CO 3 " in the 0.04 M H 2 CO 3 solution has been found by other methods to be 
 of the same order of magnitude as that of OH~. Questions. Making use of the 
 fact that the only ions in this solution are produced by the ionization of H 2 CO 3 
 (and H 2 O), what conclusion can you draw with regard to the extent to which 
 each of the reactions (/) and (^), Paragraph 1, has taken place in a solution 
 of carbonic acid? What is the approximate concentration of HCO 3 " in the 0.04 
 M H 2 CO 3 solution? List the substances present in this solution (a) at moderate 
 concentration, (b) at small concentration ( two substances), and (c) at extremely 
 small concentration (two substances). 
 
 7. The neutralization of carbonic acid. When NaOH solution is added gradu- 
 ally to a carbonic acid solution the neutralization of H 2 CO 3 takes place in steps 
 corresponding to the two steps in the dissociation of H 2 CO 3 . The main reaction 
 for the first step in the neutralization is 
 
 HoCO s + OH- + (Na + ) =* H 2 O + HCCV + (Na + ), 
 
 and the solution made from eaual molal quantities of H 2 CO 3 and NaOH is a 
 
 sodium bicarbonate soultion. The second step in the neutralization of H 2 CO 3 is 
 
 HCO 3 - + OH- == H 2 O + CO 3 - - 
 
 [39] 
 
This is the reaction for the neutralization of HCO 3 ~ by a strong base, and was 
 considered in Assignment 25. 
 
 8. Substances present in a solution of sodium carbonate. Review your notes, 
 Assignment 25, on the relative concentrations of CO 3 ~ ~, HCO 3 ~ and OH~ in a 
 0.5 M Na 2 CO 3 solution; and calculate the concentration of H + in this solution. 
 The concentration of H 2 CO 3 in this solution is also extremely small, but of course 
 is greater than the concentration of H + . Questions. How would you prepare 1 
 mol of solid Na 2 CO 3 from 1 mol of NaHCO 3 , and also from 1 mol of H 2 CO 3 ? 
 What quantity of NaOH would be required in each case? List the substances 
 which are present in a solution of Na 2 CO 3 , (a) at moderate (or high) concentra- 
 tion, (b) at relatively small concentration (two substances), and (c) at extremely 
 small concentration (two substances). 
 
 9. Substances present in a solution of sodium bicarbonate. From the fact that 
 HCO 3 ~ is a weak acid, what would you expect to be the principal substances in 
 a solution of NaHCO 3 ? Noting that the substance HCO 3 ~ is intermediate in 
 composition between H 2 CO 3 and CO 3 ~~, and considering the equilibria involved, 
 would you expect the concentration of CO 3 ~" in a solution of NaHCO 3 to be 
 greater or less than in a solution of (a) H 2 CO 3 , (b) Na,CO 3 ? The following 
 experiment, on the precipitation of CaCO 3 by NaHCO s solution at room temp- 
 erature, furnishes evidence with regard to the concentration of CO 3 ~~ in this 
 solution, (but it should be remembered that more CaCO 3 is precipitated than 
 corresponds to the actual concentration of the CO 3 ~~ since on account of the 
 displacement of equilibrium some more CO 3 ~~ forms as the CaCO 3 is precipitated. 
 Experiment. To 10 cc. M NaHCO 3 in a flask add about 50 cc. water, and 10 cc. 
 N CaCl 2 ; shake the mixture. Filter and note the amount of the CaCO 3 preci- 
 pitate. Compare the result with that obtained in the experiment with Ca ++ and 
 H,CO 3 , Paragraph 6. Repeat the experiment using 10 cc. TV (0.5 M) Na 2 CO 3 
 instead of 10 cc M NaHCO 3 . Heat the two nitrates to boiling, and test the 
 gas evolved in the NaHCO 3 experiment by passing it into a clear solution of 
 Ca(OH) 2 or Ba(OH) 2 (cf. Assignment 32, Paragraph 11). Compare the 
 amount of the precipitate now obtained with that formed in the cold solution. 
 Complete the following equation : 
 
 Ca ++ + 2HCO 3 - = CaCO 3 (solid) 
 
 and suggest an explanation for the shifting of the equilibrium when the mixture 
 is heated to boiling. The reaction that takes place when a solution of NaHCO 3 
 is heated to boiling is 
 
 2HC0 3 - = H 2 C0 3 + C0 3 " = C0 2 (gas) + H 2 O + CO 3 " 
 
 Questions. How will the concentrations of H 2 CO 3 and CO 3 ~~ in a NaHCO 3 
 solution be changed (/) when the solution is boiled, and (2) when CaCL 
 solution is added at room temperature? What conclusions can you draw with 
 regard to the relative concentrations of HCO 3 ~, and H 2 CO 3 and CO 3 " ~ in a freshly 
 prepared, cold solution of NaHCCX? 
 
 10. Experiment. Test a freshly prepared, approximately molal solution of 
 NaHCO 3 with litmus and with phenolphthalein and estimate approximately the 
 concentration of OH~ and H + in the solution. Repeat the experiment with a 
 solution which has been heated to boiling, or which has been allowed to stand 
 in the laboratory for several days. From your results in Paragraphs 9 and 10 
 state which substances are present in a pure NaHCO..- solution, (a) at moderate 
 (or high) concentration, (ft) at relatively small concentration, (two substances) 
 and (c) at still smaller concentration (two substances). 
 
 11. Summarize in tabular form the lists referred to in the last sentence of 
 each of the three Paragraphs, 6, 10 and 8, arranging- the columns in the order 
 H,CO 3 , NaHCO 3 and Na 2 CO 3 . Note the regular change in the concentration 
 of "each substance, e. g. H 2 CO 3 , when the three solutions are considered in this 
 
 [40] 
 
order. If the arrangement of the various substances in your table is not sym- 
 metrical you have probably made some mistake. Question. What is the equation 
 for the main reaction when 0.1 mol of strong acid is added to (a) 0.1 mol 
 Na,CO 3 , and (b) 0.05 mol Na 2 CO 3 ? 
 
 12. The decomposition of solid sodium bicarbonate. Experiment. Place about 
 1 gram NaHCO 3 in a hard-glass test-tube and lead the delivery tube into a 
 solution of Ba(OH) 2 , or of CaCL and NH 4 OH. Heat the tube until no more 
 gas is given off, remove the delivery tube from the solution, and allow the hard- 
 glass tube to cool. What is the evidence that CO 2 was formed? That water 
 was formed ? What must the residue be if H 2 O and CO 2 are formed in equimolal 
 quantities? (This experiment may be performed quantitatively by weighing the 
 hard-glass tube (/) empty, (2) with NaHCO 3 , and (5) with the residue.) 
 
 Experiment. Continued. Clean the delivery tube and lead it into a fresh 
 solution of Ba(OH) 2 . Pour water carefully into the hard-glass tube until it is 
 half full, but do not stir the mixture. Add 2 cc. 6 N HC1, and at once insert the 
 stopper. What is the gas evolved? Warm the mixture slowly by placing the 
 hard-glass tube in a beaker of water and heating the latter. Questions. If the gas 
 volumes had been measured, what would have been the ratio of the volumes in 
 the two parts of the experiment? If all the CO 2 evolved had been converted into 
 BaCO 3 , what would have been the relative amounts obtained in the two parts 
 of the experiment? 
 
 13. Reaction between an acid and the negative ion of a weaker acid. Give 
 examples of the action of a strong acid on a solution of the salt of a weak acid. 
 Experiment. Determine the action of dilute acetic acid on NaHCO 3 . 
 Which is the stronger acid, HAc or H 2 CO 3 ? Taking into consideration the fact 
 that HCO 3 ~, the negative ion of H 2 CO 3 , is also a weak acid (whose ions are H* 
 and CO 3 ), and the additional fact that H 2 CO 3 is a stronger acid than HCO 3 ', 
 predict what will happen when H 2 CO 3 is introduced into a solution of a car- 
 bonate. 
 
 14. Experiment. Test your answer to the last question by placing a small 
 amount of freshly precipitated CaCO 3 in about 100 cc. water and saturating the 
 mixture with CO 2 . Filter off any CaCO 3 that remains and heat the filtrate to 
 boiling. 
 
 15. Summarize all the evidence, presented in this Assignment that the reaction 
 
 2HC0 3 - = H 2 C0 3 + C0 3 " 
 
 is reversible, noting especially the experimental conditions under which the re- 
 action will proceed almost completely, (a) as written, and (b) in the reverse 
 direction. Question. Which solution would be more alkaline at room tempera- 
 ture (on account of hydrolysis), a molal solution of a salt of a weak monobasic 
 acid of the same strength as H 2 CO 3 , or a molal solution of NaHCO 3 ? Give your 
 reasoning. 
 
 16. Problems. (/) Can the following substances be present at moderate con- 
 centrations in the same solution? If not, what is formed? 
 
 (a) H 2 CO 3 and Ca ++ (/) HCO 3 - and H + 
 
 (b) H 2 CO 3 and OH' (g) HCO 3 - and OH- 
 
 (c) H 2 CO 3 and CO 3 " (h) HCCV and CO," 
 (<0 H 2 CO 3 and HCO,- (i) CO 8 and OH- 
 (e) H,CCL and H H 
 
 (2) What weight of Na 2 CO 3 can be obtained from 8.4 g. NaHCO 3 , (a) by 
 heating, and (b) by treating the bicarbonate with NaOH solution? What is 
 the least volume of N NaOH that can be used in (b) ? 
 
 [41] . 
 
(j) What volume of CO 2 at standard conditions can be obtained from 8.4 g. 
 NaHCO 3 , (a) by heating, and (b) by treating the bicarbonate with HC1 solution? 
 What is the least volume of N HC1 that can be used in (b) ? 
 
 (4) When a current of CO 2 is passed into a solution of Ca(OH) 2 a precipitate 
 is observed to form, and then dissolve. If the final solution is now heated to 
 boiling, a precipitate again appears. State what has happened in this experiment, 
 and write an equation for the main reaction in each of these stages. 
 
 ASSIGNMENT 34 
 SULPHATES, CHLORIDES AND NITRATES OF COPPER AND ZINC 
 
 1. In Assignment 34 and the four following Assignments we shall study the 
 chemistry of copper ion, Cu ++ , silver ion, Ag + , and zinc ion, Zn ++ . We shall 
 first devote our attention to the common soluble salts, and to the method of 
 transforming one salt into another ; and then we shall study the sparingly soluble 
 compounds, the methods of dissolving them, and the equilibria involved. The 
 same method of treatment will be used later in studying the chemistry of ions 
 of other metals, and many of the results now obtained in studying copper, silver 
 and zinc are also true for other metals. Read again Paragraph 2, Assignment 31. 
 
 2. Solubility of nitrates, acetates, chlorides and sulfates of metals. The 
 nitrates and acetates of all metals are soluble in water and this is also true foi 
 nearly all the chlorides and sulfates. Question. Give an example of a diffi- 
 cultly soluble chloride and of a difficultly soluble sulfate. Give the formulas of 
 the nitrates, acetates, chlorides and sulfates of copper, silver and zinc. 
 
 3. Preparation of sulfates from nitrates and chlorides. Experiment. Test 
 the relative volatility of HC1, HNO 3 and H 2 SO 4 by evaporating a few drops of 
 a concentrated solution of each of these acid's in a casserole (out of doors or in 
 a fume closet). Cf. Hildebrand, pages 173-175. Suggest a method of preparing 
 solid copper sulfate from copper nitrate based on the difference in volatility 
 of HNO 3 and H 2 SO 4 . Experiment. Try your method with 15 cc. of the labora- 
 tory solution of copper nitrate. Recrystallize the copper sulfate by mixing the 
 residue, after it is cold, with 2 or 3 cc. water, heating the mixture to boiling 
 and letting it cool slowly. Wash the crystals with a very little water, test a 
 small portion for NO 3 ", and save the remainder. 
 
 4. Suggest a similar method for the preparation of solid zinc sulfate from 
 zinc chloride. How would you test whether the final product is free from 
 chloride ? 
 
 5. The problem of preparing a soluble chloride or nitrate from a soluble 
 sulfate will be considered in the following Assignment. 
 
 6. Conversion of soluble chlorides into nitrates, and nitrates into chlorides. 
 Experiment. Mix 2 cc. concentrated HNO 3 solution and 6 cc. concentrated 
 HC1 solution, and let the mixture stand. Evaporate a few drops of the solution 
 to dryness (out of doors or in a fume closet) and note if there is a residue. 
 The gradual deepening of the color of the solution at room temperature proves 
 not only that a reaction is taking place but that this reaction is not a rapid one. 
 Questions. How would the speed of the reaction be altered (a) by raising the 
 temperature, and (b) by using less concentrated acids? Note. Although this 
 reaction does not take place in a dilute solution which contains H + , Cl~ and 
 NO 3 ", it can be made to do so by concentrating the solution by evaporation to 
 a small volume. Since both Cl~, and NO 3 ~ are used up in this reaction and the 
 products are volatile, we can make use of this reaction (/) to remove Cl~ from 
 a solution by heating with excess of concentrated HNO 3 , or (2) to remove 
 
 [42] 
 
NO 3 ~ from a solution by heating with excess of concentrated HC1. Question. 
 If excess of concentrated HNO 3 is added to a small amount of NaCl and the 
 mixture evaporated just to dryness, what would you expect the solid residue 
 to be. Note. The mixture of concentrated HNO 3 and HC1 is called aqua regia, 
 and the reaction that takes place is NO 8 - + 3 Cl~ + 4H + = NOC1 + CL, + H 2 O. 
 
 7. Experiment. To about 0.2 g. copper chloride (or 2 cc, N ZnCl 2 solution) 
 in a casserole, add 3 cc. concentrated HNO 3 and stir; add one or two drops of 
 this mixture to 10 cc. water, test this solution for Cl~ and set it aside for com- 
 parison with later results. Evaporate the mixture of the chloride and con- 
 centrated HNO 3 just to dryness (out of doors or in a fume closet)., add 5 cc. 
 concentrated HNO 3 being careful to dissolve any solid on the side of the 
 dish, and test one or two drops of the solution for Cl~ as before. Again 
 evaporate the mixture just to dryness, and test a portion of the residue for 
 chloride. Questions. What is the final solid residue? What conclusion can 
 your draw from a comparison of the amounts of the AgCl precipitates in your 
 three tests? Does this experiment furnish any evidence that the reaction is 
 slow, even at a temperature in the neighborhood of 100, when the concentration 
 of one of the reacting substances is small? 
 
 8. How would you convert copper nitrate, or zinc nitrate, into the chloride? 
 
 9. Treatment of silver chloride with hot concentrated nitric acid and sul- 
 furic acid. Experiment. Prepare some silver chloride, and wash it with dilute 
 HC1 and then with water. Place a small amount of the silver chloride in a 
 porcelain dish, and 3 or 4 cc. concentrated HNO 3 , evaporate the mixture to 
 dryness. Add a little water and test the solution for Ag + . Repeat the experi- 
 ment with the residue. Compare the result with that in the experiment in 
 Paragraph 7, and state how the speed of the reaction between Cl~ and NO 3 ~ in 
 acid solution depends on the concentration of Cl~. 
 
 10. Test the action of hot concentrated H 2 SO 4 on silver chloride by heating 
 for several minutes a small amount of the salt with 3 or 4 cc. of the acid in a 
 porcelain dish covered with a watch glass. Set the dish aside until it is cool, 
 pour the mixture into water, and test the solution for Ag + . Note. The chloride 
 is expelled as HC1 gas in this experiment, but since the temperature may have 
 exceeded 300 it is not surprising that the result is different from that obtained 
 in the experiment with AgCl and concentrated HNO 3 . 
 
 11. Problems. (/) What will be the solid residue when each of the follow- 
 ing solutions is evaporated to dryness? 
 
 (a) A solution which contains Ag + , H + and NO 3 ~. 
 
 (b) A solution which contains Cu ++ and NO 3 ~ at small concentrations and 
 H + and SO 4 '~ at large concentrations. 
 
 (c) A solution which contains Cu ++ and NO 3 ~ at large concentrations and 
 H + and SO 4 ~~ at small concentrations. 
 
 (d) A solution of zinc sulfate to which nitric acid has been added. 
 
 (2) How would you prepare solid silver sulfate from (a) solid silver nitrate, 
 (b) solid silver acetate? 
 
 (j) When solid copper sulfate is prepared by crystallization from an aaueous 
 solution it contains water of crystallisation and its formula is CuSO 4 5H 2 O. 
 When this blue substance is heated, white anhydrous CuSO 4 is finally obtained. 
 Write the equation. The reverse reaction takes place when air saturated with 
 water vapor is passed over the anhydrous salt at room temperature. If a vessel 
 containing CuSO 4 5H 2 O were evacuated for a few -minutes and then closed, 
 what substances would be present in the closed vessel? What reason do you 
 have for the conclusion that the pressure of water vapor in the vessel at room 
 temperature must be less than the partial pressure of water vapor .at the same 
 
 F431 
 
temperature? How would the pressure in the vessel be altered by raising the 
 temperature? Give two other examples of a reversible reaction in which one 
 solid substance dissociates into another solid and a gas, and in each case give 
 an approximate value for the equilibrium pressure at some definite temperature. 
 (Refer to the lectures on the dissociation of calcium carbonate and calcium 
 hydroxide when heated.) 
 
 ASSIGNMENT 35 
 HYDROXIDES OF COPPER, SILVER AND ZINC 
 
 References. Hildebrand, Chapter V pages 72-73. 80-82; Chapter XIII pages 
 
 194-195. 
 
 1. We have learned from the laboratory work and the lectures that, the 
 hydroxides of sodium, potassium and the other alkali metals, and ammonium 
 hydroxide are readily soluble in water, and that the hydroxides of barium, 
 strontium and calcium are moderately soluble (with the solubility decreasing 
 rather rapidly in the order named). The hydroxides of all other metals are 
 difficulty soluble in water. 
 
 2. Cuprlc hydroxide. State what solutions you would mix to prepare 
 Cu(OH) 2 . Point out the substances involved in the equilibrium between the 
 solid and its saturated solution and predict what will happen when the solid is 
 treated with HC1, HNO 3 or H 2 SO 4 solution. Experiment. Test you* 
 answer by preparing some copper hydi oxide and treating it with acids. Also 
 determine if it dissolves in 0.5 'N NaOH. 
 
 3. Cuprlc oxide. Experiment. Collect some Cu(OH) 2 on a filter and 
 wash it once with water. Mix some of the Cu(OH) 2 with water and heat the 
 mixture to boiling. Place the remainder of the Cu(OH) 2 in a procelain dish 
 and heat gently. Cupric oxide has formed. Does the reverse reaction between 
 CuO and H 2 O take place at room temperature? Predict what will happen when 
 NaOH solution is added to a solution of Cu(NO 3 ) 2 at 100, and test your 
 answer. Test the action of HC1, HNO 3 or H 2 SO 4 solution on CuO. Question. 
 How would you prepare solid CuCl 2 from Cu(OH) 2 or CuO? 
 
 4. Sliver oxide. Experiment. To 10 cc. 0.1 TV AgNO 3 solution add 
 NaOH solution slowly until, after shaking, the solution reacts strongly alkaline 
 to litmus. Silver oxide, Ag 2 O, is formed. Filter, acidify the filtrate with 
 HNO 3 and test it for Ag + . Write the equation for the reaction between Ag + 
 and OH~ to form Ag 2 O. Point out the substances involved in the equilibrium 
 between the solid substance and its saturated solution, and predict what will 
 happen when the solid substance is treated with HNO 3 solution. Test your 
 answer by an experiment. 
 
 5. Relation between oxides and hydroxides of metals. Contrast the action of 
 water or water vapor at room temperature on calcium oxide, and on copper 
 oxide or silver oxide. Review Problem 3, Assignment 34; and note that the 
 reaction 
 
 Oxide of a metal (solid) + H 2 O (gas) = Hydroxide of a metal (solid) 
 takes place completely in the case of oxides of the alkali metals, as Na 2 O, does 
 not have any tendency to take place in the case of oxides of the more noble 
 metals, as Ag 2 O, and can be shown to be reversible for many oxides which 
 are intermediate between these two extremes. (In many cases equilibrium is 
 reached only very slowly,' especially at low temperatures.) On account of this 
 relation between oxides and hydroxides of metals, it is not surprising that oxides 
 of metals in general are capable of neutralizing acids. Questions. What is 
 anhydride of a base? Give examples. Is it always correct to say than an 
 
 [441 
 
anhydride of a base will react with water to form the base? What is an acid 
 anhydride!' Give examples. Also give examples of reactions between an acid 
 anhydride and a base, and between an acid anhydride and a basic anhydride. 
 
 6. Acids which contain oxygen may be considered to be related to the hydrox- 
 ides of non-metals. Thus H 2 SO 4 may be regarded as H 6 SO 6 or S(OH) from 
 which 2 molecules of water have been withdrawn 
 
 S(OH) 6 = H 2 SO 4 + 2 H 2 O; 
 
 and the formula of phosphorous acid, H 3 PO 3 , may be written P(OH) 3 . The 
 characteristic properties of these acids, to yield negative ions containing oxygen 
 on dissociation or neutralization, is then to be attributed to the existence of 
 strong bonds between the non-metal and the oxygen in the compound; which is 
 really due to the great tendency of non-metals to hold electrons firmly. On the 
 other hand metals in compounds do not hold electrons firmly, and positive ions 
 of metals are formed when bases or salts dissociate. However, certain elements 
 which form a positive ion also can form a negative ion containing oxygen; the 
 hydroxide of such an element has the properties of a base in that it can neutralize 
 an acid, and also has the properties of an acid since it can neutralize a base: it 
 is said to be amphoteric. 
 
 7. Zinc hydroxide. Experiment. Prepare some Zn(OH) 2 by treating a 
 solution containing Zn ++ with a very small amount of NaOH. Collect the 
 Zn(OH) 2 on a filter, treat a portion with a strong acid, and another portion with 
 NaOH or KOH solution. Question. What evidence is furnished by this 
 experiment that zinc hydroxide is an amphoteric substance? The reaction be- 
 tween zinc hydroxide and excess OH~ is analogous to the first stage in the neu- 
 tralization of carbonic acid : 
 
 H ZnOo (solid) + OH- = H,O + HZnCX- 
 H 2 CO S + OH- H 2 O + HCO S - 
 
 The negative ion, HZnO 3 ", is usually zincate ion, though strictly speaking this 
 name should be reserved for the ion ZnO 2 ~~. The latter substance probably 
 forms to some extent when the concentration of OH" is very large. Solid 
 NaHZnO, has not been prepared, but Na^ZnOo can be made by fusing zinc oxide 
 with NaOH. 
 
 8. Predict what will happen when a solution of a strong acid is added slowly 
 to sodium zincate solution. Test your answer by an experiment. 
 
 9. Preparation of a nitrate or chloride from a soluble sulfate. Experiment. 
 Prepare a solution of copper nitrate from copper sulfate solution by precipitat- 
 ing copper hydroxide from a very dilute CuSO 4 solution, washing the Cu(OH) 2 
 precipitate, and transforming it into the nitrate. Test the product for sulfate. 
 
 10. Suggest a second method based on the removal of the sulfate ion by 
 means of barium ion (cf. Hildebrand. pages 207-208). Try your method, and 
 test the copper nitrate solution for SO 4 ~~ and for Ba ++ . Which is the better 
 method to use on a small scale in the laboratory? 
 
 11. Suggest two methods of preparing solid zinc chloride from zinc sulfate. 
 
 12. Problems. (/) Write equations for the action of HC1 solution in excess 
 on calcium oxide, on ferrous and ferric oxides (FeO and Fe.X) 3 ), on ferrous and 
 ferric hydroxides, and on sodium zincate solution. 
 
 (2) A solution containing 0.2 mol Na 2 SO 4 is mixed w r ith a solution contain- 
 ing 0.1 mols Ba(NO 3 ) 2 , the precipitate is removed by filtration, the wash water 
 is run into the filtrate until the final volume of the latter is 1 liter. What sub- 
 stances are present in the final solution and what is the concentration of each? 
 Suggest a method of preparing pure sodium nitrate from sodium sulfate. 
 
 ( j) Give two methods of preparing CuCL from CuSO 4 . 
 
 [45] 
 
(4) Give an example of an amphoteric hydroxide other than Zn(OH) 2 , and 
 write the equation for its reaction with (a) a strong acid, (b) a strong base. 
 
 ASSIGNMENT 36 
 COMPLEX IONS OF COPPER, SILVER AND ZINC WITH AMMONIA 
 
 References. Hildebrand, pages 184 and 202-205. 
 
 1. The ions of certain metals have the power of forming compounds with 
 NH 3 , which are examples of "complex ions". The ammonia is supplied by 
 adding NH 4 OH solution. There is an equilibrium between the complex ion, 
 NH 3 or NH 4 OH, and the ion of the metal; and when the NH 4 OH is present 
 in excess the concentration of the ion of the metal is often very small. Some 
 examples will be considered in the present Assignment. 
 
 2. Experiment. Dilute 5 cc. N CuSO 4 with 20 cc. water and add two or 
 three drops 6 N NH 4 OH. The precipitate is Cu(OH) 2 . Question. What 
 evidence with regard to the solubility of Cu(OH) 2 in water is furnished by this 
 experiment? (Consider the low concentration of OH~ in NH 4 OH solution and 
 the result of the experiment in Paragraph 6, Assignment 32. 
 
 3. Continue to add the NH 4 OH solution until, after shaking, the solution is 
 clear. Observe the volume of NH 4 OH used. What is the evidence that a new 
 substance is formed? Repeat the experiment with the same volume oi 6 N 
 NaOH instead oi 6 N NH 4 OH. State the reasoning by which you may con- 
 clude that the substance obtained in the experiment with NH 4 OH solution is 
 not formed by a reaction between OH~ and Cu(OH) 2 similar to that between 
 OH~ and Zn(OH) 2 . Try an experiment with CuSO 4 solution and NH 4 C1 
 in order to determine if Cu ++ reacts with NH 4 + in the same way as with NH 3 
 or NH 4 OH. 
 
 4. The new substance formed when Cu(OH) 2 dissolves in NH 4 OH solution 
 is Cu(NH 3 ) 4 ++ . The final solution contains NH 4 OH, and OH- must be present 
 at an appreciable concentration ; it may be concluded that the concentration of 
 Cu ++ is extremely small, since, otherwise, Cu(OH) 2 would precipitate. Questions. 
 Which solution has the greater concentration of Cu ++ , a saturated solution of 
 Cu(OH) 2 or a solution of the complex ion containing NH 4 OH? How is the 
 equilibrium between Cu(OH) 2 and its saturated solution affected when NH.,OH 
 is added? 
 
 5. The equation for the main reaction when a solution of a copper salt is 
 treated with excess NH 4 OH is 
 
 Cu ++ +4NH 4 OH = Cu(NH 3 ) 4 ++ + 4H 2 O. 
 
 A simple method of demonstrating that the reverse reaction will take place is to 
 remove the NH 4 OH. Suggest two methods of doing this and try your methods. 
 
 6. Zinc and silver also form complex ions with ammonia, 
 
 Zn(NH 8 ) 4 " and Ag(NH 3 ) 2 + . 
 
 Experiment. To a dilute solution of ZnSO 4 add NH 4 OH solution drop by 
 drop, shaking the mixture after each addition of NH 4 OH; when some but not 
 all of the Zn(OH) 2 has dissolved test the solution with trinitrobenzol to determine 
 approximately the OH~ concentration. Repeat the experiment, using NaOH 
 instead of NH 4 OH, and note again the OH- concentration in the solution when 
 Zn(OH) 2 is partly but not completely dissolved. Question. From a comparison 
 of the OH~ concentration at equilibrium in the two cases what conclusion can 
 you draw with regard to the presence of zinc as a different substance in the 
 NH 4 OH and NaOH solutions? 
 
 [46] 
 
7. Experiment. Prepare some silver oxide and some silver chloride, and 
 treat each substance with NH 4 OH solution. Predict what will happen when 
 the resulting solutions are acidified with nitric acid, and test your answers by 
 experiments. 
 
 8. Which solution would contain a greater concentration of Ag + , a saturated 
 solution of silver acetate (see Assignment 24, Paragraph 5) or a solution which 
 contains the complex ion and NH 4 OH? What then must be the action of 
 NH 4 OH solution on solid AgAc? 
 
 9. Questions. When AgCl is dissolved in excess NH 4 OH solution, what is the 
 principal positive ion, and the principal negative ion, in the resulting solution? 
 When o.i mols AgCl are dissolved and the final volume is 100 cc. what is the 
 approximate concentration of each of these ions? Is the concentration of Ag* 
 smaller or larger than in a saturated solution of AgCl in water? 
 
 10. When Cu(OH) 2 is dissolved in excess NH 4 OH solution what is the 
 principal negative ion in the final solution? , It is important to realize that 
 Cu(NH 3 ) 4 (OH) 2 , Zn(NH 3 ) 4 (OH) 2 and Ag(NH 3 ) 2 OH are soluble, strong 
 bases. 
 
 11. We shall next consider the action of NH 4 OH solution on silver iodide, a 
 salt which is much less soluble in water than either silver oxide or silver chloride. 
 Solid Agl, suspended in water, is in equilibrium with Ag + and I~ in the 
 saturated solution. When NH 4 OH is added some Ag + must react with it to 
 form Ag(NH 3 ) 2 + , and some solid Agl must dissolve to establish again the 
 solubility equilibrium. We can therefore predict that Agl must be more soluble 
 in NH 4 OH solution than in water, but it is not safe to predict that a large 
 amount of it will dissolve. Experiment. Prepare a solution of silver oxide 
 or silver chloride in NH 4 OH and add to this solution a few drops of 0.1 N potas- 
 sium iodide solution. In another experiment collect some solid silver iodide on 
 a filter paper, wash it with water, and treat it with dilute NH 4 OH solution. 
 Devise and try an experiment with the solution thus prepared to determine 
 whether an appreciable amount of Agl has dissolved. 
 
 12. Predict what will happen when silver chloride is treated with a solution 
 of potassium iodide. Try the experiment. 
 
 13. The ions of the alkali metals, as Na + , K + , etc.; those of the alkaline earth 
 metals, as Ca ++ , Sr TT and Ba ++ ; and many other positive ions do not combine 
 with NH 3 to form complex ions. Question. Will a difficultly soluble salt of 
 calcium, as CaCO, or CaC 2 O 4 , dissolve in NH 4 OH solution to a greater extent 
 than in water? 
 
 14. Problems, (i) Write equations for the main reactions in the following 
 experiments : 
 
 (a) A few drops of NH 4 OH solution are added to a solution contain- 
 ing Zn ++ . 
 
 (b) Excess NH 4 OH solution is added to a solution containing Zn ++ . 
 
 (c) An ammoniacal solution of silver chloride is acidified with nitric acid. 
 
 (2) The solubilities of AgCl and Agl in water are 10~ 4 and 10~ 7 mols per 
 liter, respectively. In two experiments with NH 4 OH solution each substance 
 is found to be one thousand times more soluble than in water. How many grams 
 of each has dissolved in i liter of the NH 4 OH solutions? 
 
 (j) Give two examples of complex ions other than those with ammonia. 
 
 (4) From the positions of Cu, Ag, and Zn in the Periodic System (Hildebrand 
 page 257), state- what other ions might be expected to form ammonia complexes. 
 Check your answer by referring to Hildebrand, page 269. 
 
 [47] 
 
ASSIGNMENT 37 
 CARBONATES AND SULFIDES OF COPPER, SILVER AND ZINC 
 
 References. Hildebrand, pages 199-201 and 195-197. 
 
 1. In Assignment 33 we learned that when carbon dioxide dissolves in 
 water the solution contains the weak acid H 2 CO 3 , and in the saturated solution 
 we recognized an equilibrium between H 2 CO 3 in the solution and CO 2 in the 
 gas space. We found that carbonic acid ionizes in two steps and can be neu- 
 tralized in stages to give solutions of the two corresponding types of salts, and 
 that when these salts are treated with a strong acid carbon dioxide is evolved. 
 In the present Assignment we shall continue the study of difficultly soluble car- 
 bonates, and we shall deal with the salts of another dibasic acid, H 2 S, which is 
 in many respects analogous to carbonic acid. 
 
 2. Hydrogen sulfide, like carbon dioxide, is a gas which is soluble in water, 
 and when the solution is saturated an equilibrium is established between gaseous 
 and dissolved H 2 S. The solution has the properties of a weak acid which ionizes 
 in two stages as does carbonic acid. Sulfides, in general, are less soluble than 
 carbonates ; in the case of the alkali metals both the sulfides and carbonates are 
 readily soluble, while in the case of the alkali earth metals the sulfides are 
 readily soluble and the carbonates are not. Questions. Write the two equations 
 which show the ionization of H 2 S (dissolved in water) in two stages. What 
 are the names of the negative ions formed ? In analogy with the ions of car- 
 bonic acid, which is the stronger acid, H 2 S or HS"? What are the principal 
 substances present in a solution of sodium sulfide, Na 2 S? In a solution of 
 sodium hydrogen sulfide, NaHS? What reactions take place when a strong acid 
 is added to any solution containing sulfide ion? How, then, will the solubility of 
 difficultly soluble sulfides be affected by the addition of a strong acid? 
 
 3. Experiment. Try the action of Na 2 CO 3 solution on 0.1 N solutions of 
 Cu+ + , Ag + , and Zn ++ . In each case collect the precipitate on a filter, wash it 
 with water, and test a portion for carbonate. Predict the action of nitric acid 
 solution and of ammonium hydroxide on these precipitates, and test your pre- 
 diction by experiments. 
 
 4. In the above experiment with Cu ++ and in many other cases, the precipitate 
 obtained by the addition of Na 2 CO 3 to a solution of a salt is not a pure car- 
 bonate, but a substance which contains both the carbonate and hydroxide 
 radicals. Such substances, called basic carbonates, are difficult to obtain pure, 
 the actual precipitates consisting of variable mixtures of these basic carbonates, 
 carbonates, and hydroxides. 
 
 5. Experiment. Prepare 20 cc. of 0.1 N solutions of Cu ++ , Ag% and 
 Zn + +. To each add I cc. 6 N H 2 SO 4 and pass in H 2 S gas* until the liquid is 
 saturated. To determine this, close the end of the test-tube or flask, shake 
 thoroughly and test the odor cautiously. Collect the precipitates on separate 
 filters. In the experiment with Zn ++ and H 2 S, to the filtrate (or to the clear 
 solution if there was no precipitate), add 10 cc normal NaAc solution, and 
 again saturate with H 2 S gas. Question. What effect has the addition of Ac" on 
 the concentration of H + ? 
 
 6. H 2 S and HS~ are weaker acids than H 2 CO 3 and HCO 3 ~, respectively. What 
 can you conclude about the relative concentrations of S~~ and CO. { ~~ in solutions 
 of H 2 S and H 2 CO 3 ? It will be recalled that 'H 2 CO 3 does not precipitate CaCO 3 
 
 * Caution. H 2 S is poisonous. Work out of doors if possible and do not breathe the gas. 
 
 [481 
 
from a solution of CaCL (Assignment 33, Paragraph 6). Question. What 
 conclusion can you draw with regard to the relative solubility of CaCO 3 and 
 the sulfides prepared in the preceding Paragraph? 
 
 7. From the equation for the main reaction 
 
 M ++ + H 2 S = MS (solid) + 2H+ 
 
 it is evident that both M ++ and H* are competing for S~~ and that an increase 
 in the H + concentration tends to reverse the reaction, as we have seen in the 
 case of ZnS. However in the case of a very insoluble sulfide, CuS for example, 
 even though H + concentration is fairly large the concentration of M ++ will be 
 negligible if the solution is saturated with H 2 S. Determine by experiment which 
 of these sulfides will dissolve completely in 2 N H 2 SO 4 . Which then has the 
 greater solubility in water? Write the two equations which show the mechanism 
 of the main reaction. State how you would precipitate a sulfide which is spar- 
 ingly soluble in water, but many times more soluble than ZnS (cf. the precipita- 
 tion of CaCO 3 ). 
 
 8. Experiment. To solutions containing Cu(NH 3 ) 4 ++ , Ag(NH 3 )o + , and 
 Zn ( NH 3 ) 4 + +, in separate experiments, add (NH 4 ) 2 S solution. Repeat the 
 experiments, using H 2 S gas instead of (NH 4 ) 2 S solution. State in each case 
 which has the smaller concentration of the ion of the metal, a solution con- 
 taining the complex ion, or a solution saturated with the sulfide. State also 
 which is the less soluble in water, Cu(OH) 2 or CuS, Ag 2 O or Ag 2 S, Zn(OH) 2 
 or ZnS. From the fact that the sulfides are less soluble than the carbonates, 
 predict the action of H 2 S on mixtures of basic copper carbonate and water, 
 silver carbonate and water, and zinc carbonate and water. Test your answers 
 by experiment. 
 
 9. Experiment. Test the action of H 2 S on solutions of nitric acid of 
 different concentrations, 0.1 N, 2.0 N, and 6 N. Saturate each solution with 
 H 2 S gas and heat the mixture almost to boiling. The white precipitate formed 
 is solid sulfur and the principal reaction is 
 
 3H 2 S + 2NO 3 - + 2H + = 3S (solid) + 2NO + 4ELO. 
 
 Questions. What evidence is furnished by your experiments that this is a slow 
 reaction? How is the speed affected by the concentration of the nitric acid, and 
 by the temperature? (Note. The products of the reaction, S, NO, and H 2 O, 
 do not react with each other under any ordinary conditions of temperature and 
 pressure, and it is practically impossible to realize an equilibrium involving the 
 substances shown in the equation.) Because of this reaction nitric acid in a 
 strongly acid solution will remove sulfide ion from the solution more completely 
 than will H + alone since it destroys the H 2 S which is in equilibrium with sulfide 
 ion. We may therefore expect that nitric acid will be a better solvent for 
 difficulty soluble sulfides than sulfuric or hydrochloric acid. 
 
 10. Experiment. Collect some CuS on a filter, transfer it to a casserole, 
 add 2 normal HNO 3 and boil the mixture. Filter, and test the filtrate for 
 Cu ++ . The residue collected on the filter is sulfur; the dark color is due to a 
 little CuS enclosed in the sulfur. Complete the equation 
 
 3CuS (solid) + 2NO 3 - = 3S (solid) + 2NO 
 
 Problems. (/) Hydrogen sulfide is passed into 1 N CuSO 4 until the precipita- 
 tion of copper sulfide is complete. The liquid is filtered and H 2 S is removed by 
 boiling. What concentration of H + is left in the solution. 
 
 (2) Write the equations for the neutralization of (a) one mol H.,S with one mol 
 OH-, (b) H 2 S with excess OH~, (c) HS~ with OH~. 
 
 (j) Predict the effect of NH 4 OH on the solubilities of CuS, ZnS and 
 
 T491 
 
Ag 2 S. Give your reasoning. (If you try the experiments use freshly prepared 
 sulfides moistened with water which contains H 2 S.) 
 
 (4) Give a method for the preparation ot 
 
 (a) CuSO 4 solution from CuS. 
 
 (b) Ag 2 S(solid) from AgCl. 
 
 (c) ZnCO 3 (solid) from ZnS. 
 
 (d) Na 2 CO 3 solution from Na 2 S solution; see experiment at end of 
 Paragraph 8, and Hildebrand, pages 207-208. 
 
 ASSIGNMENT 38 
 REVIEW OF THE CHEMISTRY OF THE POSITIVE IONS ALREADY CONSIDERED 
 
 1. This Assignment is inserted in order that the student may summarize and 
 study the chemistry of the positive ions considered in the last seven Assign- 
 ments. If the student makes an accurate summary, and notes carefully similar- 
 ities and differences in the results for the various ions, he will find that it is not 
 difficult to remember the large number of facts involved. In later work, as 
 soon as the chemistry of another ion has been investigated, the results should 
 be summarized, and compared with the results for the ions already studied. 
 
 2. Summary of results. Tabulate acording to the following plan the results 
 which you have obtained in earlier Assignments ; refer to your laboratory notes 
 whenever necessary. Write in a horizontal line across a double page of your 
 notebook the names of the positive ions, silver, cupric copper, zinc, calcium, 
 sodium, potassium and ammonium ; in the vertical columns write the following : 
 
 (/) the formulas of the ions, including complex ions; 
 (2) the formulas of the compounds which are readily soluble in water, 
 (j) the formulas of the compounds which are only moderately soluble, as 
 Ca(OH) 2 , CaS0 4 , Ag 2 SO 4 . 
 
 (4) the formulas of the insoluble substances (the term difficultly soluble 
 or sparingly soluble is preferable). Indicate which is the least soluble com- 
 pound of each metal. When any ion or difficultly soluble compound is colored, 
 give the color in the table. Next make as general statements as you can about 
 the solubilities in water of the nitrates, chlorides, sulfates, carbonates, hydroxides 
 (or oxides), and sulfides of these and other metals. 
 
 3. Show how you can apply your knowledge of the equilibria involved, in 
 choosing reagents to dissolve the various difficultly soluble substances listed in 
 the preceding Paragraph. In this connection, for each complex ion compare 
 the concentration of the ion of the metal in a solution which contains the complex 
 ion with that in a saturated solution of the various difficultly soluble salts of 
 the metal; cf. Hildebrand's discussion of the reactions of silver ion, pages 
 204-205, and the following example. Since ZnCO 3 dissolves readily in excess 
 NH 4 OH solution, and ZnS does not, it may be concluded that in a solution 
 which contains Zn (NH 3 ) 4 ++ and NH 4 OH the concentration of Zn ++ is less than 
 in a saturated solution of ZnCO 3 but greater than in a saturated solution of ZnS. 
 
 4. The Summary in Paragraph 2 was obtained from the results of experiments 
 on the reactions between the positive ions and various reagents. Conversely 
 the results of such experiments can be predicted by means of the Summary and 
 a knowledge of the various equilibria involved. 
 
 [50] 
 
5. Table of Reactions. Prepare for future reference a table of the reactions 
 between dilute solutions of the positive ions considered above and the following 
 reagents : 
 
 (a) HC1 or a soluble chloride, 
 
 (b) NaOH in small amount, 
 
 (c) NaOH in excess, 
 
 (d) NH 4 OH in small amount, 
 
 (e) NH 4 OH in excess, 
 (/) H 2 C0 3 or CO, gas, 
 (g] a soluble carbonate, 
 
 (A) H,S in 0.3 N H + solution, 
 
 (*) a soluble sulfide as Na 2 S, or (NH 4 ) 2 S in NH. t OH solution, 
 
 (/ ) sodium cobaltinitrite in the presence of a few drops of HAc. 
 
 When no reaction takes place mark X at the proper place in the table. In 
 other cases write the formulas of the substances formed, note the colors, and 
 mark the precipitates by underlining the formulas. Be sure that you can 
 write the equation for each reaction. Try experiments whenever it is neces- 
 sary, e. g. when your laboratory notes are incomplete or seem to be incorrect, 
 and when there is a blank space in your table. 
 
 6. Test your knowledge of the reactions just considered by writing out from 
 memory portions of the Table of Reactions, Paragraph 5 ; write equations for the 
 reactions that take place, and discuss briefly the equilibria involved. 
 
 Note. It is recommended that the student proceed with Assignment 51, the 
 first in the Section entitled Systematic Qualitative Analysis. It is important that 
 the student become thoroughly familiar with the reactions already considered 
 before proceeding to study other reactions, and this Assignment deals only with 
 the ions already considered. 
 
 51 
 
SECTION IV 
 REACTIONS OF IONS, CONTINUED 
 
 ASSIGNMENT 41 
 OXIDATION AND REDUCTION. REPLACEMENT REACTIONS. ELECTRICAL CELLS 
 
 References. Hildebrand, Chapter 'VI pages 93-94, Chapter V page 70, Chapter 
 
 XV pages 234-237. 
 
 1. In Assignment 41 and the three following Assignments we shall consider 
 examples of oxidation and reduction reactions. In these reactions changes of 
 valence occur, and they are therefore easily distinguished from other reactions 
 such as the formation of weak electrolytes or the precipitation of difficultly soluble 
 salts. The following reactions can now be recognized as oxidation and reduc- 
 tion reactions : 
 
 Zn (solid) + 2H + = Zn ++ + Ho (Assignment 3) 
 
 H, + CuO (solid) = Cu (solid) + HoO (Assignment 4) 
 
 Ca (solid) + 2H 2 O = Ca +V + 2OH- + H, (Assignment 32) 
 
 3H 2 S + 2NO 3 - + 2H + == 3S + 2NO + 4EUO (Assignment 37) 
 
 3d- + NO 3 - + 4H + = 3/2 C1 2 + NO + 2H.,O 
 
 = C1 2 + NOC1 + 2H 2 6 (Assignment 34) 
 
 Point out the changes of valence in each reaction. 
 
 2. In the first of the above reactions the oxidation of Zn to Zn ++ and the 
 reduction of 2H + to H 2 consist simply in the transfer of two electrical charges 
 from two hydrogen atoms to one zinc atom (which in reality is the transfer 
 of two electrons from the zinc atom to the two hydrogen atoms). In the pres- 
 ent Assignment we shall study a number of examples which are characterized 
 by such an interchange of electrical charges. The reaction between copper and 
 silver ion will first be studied quantitatively and then a number of other cases 
 will be taken up qualitatively. 
 
 3. Experiment. Obtain at the office 100 cc. 0.10 N silver nitrate, and a 
 special filter paper which burns to give an ash of negligible weight. Clean a 
 straight piece of thin copper wire and weight about 0.3 g. to 5 mg. Cover the 
 copper wire with 100 cc. of 0.10 N silver nitrate solution and let it stand until 
 the copper has disappeared. The precipitate formed is metallic silver. Questions. 
 What substance is responsible for the color of the final solution? What is 
 the color of silver ion? Heat a porcelain crucible, let it cool, and. weight to 5 
 mg. Collect the silver by filtering through the special filter paper and wash 
 the precipitate until the washings are colorless. (Save the filtrate, and place in 
 it another piece of copper in order to recover the rest of the silver.) Fold the 
 filter paper so that it will fit into the crucible, heat with a small flame until it is 
 dry, and then with a larger flame until the paper is completely consumed. Let 
 the crucible cool and again weigh to 5 mg. Place all the silver in the jar marked 
 "Silver Waste." 
 
 4. Calculations. Calculate (a) the weight of silver formed in the experi- 
 ment; (b) the weight of silver precipitated by one gram atom of copper. Com- 
 pare the latter number with the atomic weight of silver. (The formula of 
 silver ion is Ag + .) What is the formula of copper ion? 
 
 5. Questions. Give the formulas of silver nitrate, silver sulfate, copper 
 nitrate and copper sulfate. If one gram molecule of silver ion has associated 
 with it 96,500 coulombs of positive electricity what quantity of electricity is 
 
 [521 
 
associated with one gram molecule of copper ion? With a gram molecule of 
 nitrate ion? Write the equation for the action of silver ion on copper. How 
 would the result of this experiment have been influenced by replacing the silver 
 nitrate with silver sulfate solution? 
 
 6. Experiment. Into 10 cc. N Cu(NO 3 ) 2 place a piece of lead. 
 
 7. Repeat the experiment with a solution of lead nitrate, Pb(NO 3 ) 2 , and 
 zinc. 
 
 8. Write the equations for all the reactions between metals and positive ions 
 which you have studied (Assignments 3 and 41). Questions. The negative 
 ion present in each experiment is not shown in the equation; what experi- 
 mental fact is thereby implied? If in one of the above experiments the dry 
 salt had been used instead of its solution, would the reaction occur sufficiently 
 rapidly at room temperature to be detected? Try an exepriment, if necessary, 
 with lead and solid copper sulfate. 
 
 9. Experiment. Determine by experiment which of the metals used in 
 this Assignment will dissolve readily in dilute hydrochloric acid (or sulfuric 
 acid).* 
 
 10. Arrange the metals used in these experiments in the order in which 
 they will replace each other from solution. After referring to your lecture 
 notes state where hydrogen belongs in this replacement series. From this 
 replacement series predict what will happen when, 
 
 Silver nitrate solution is treated with lead. 
 Copper nitrate zinc. 
 
 Copper nitrate silver. 
 
 Experiment. Test the correctness of your prediction. 
 
 11. Replacement series for the halogens. Experiment. Make a small gas 
 generator (see Assignment 33, Paragraph 3) and charge the flask with about 
 2 g. of manganese dioxide. Pour 5 cc. concentrated HC1 into the flask 
 and warm gently until a continuous stream of chlorine is evolved. Pass a few 
 bubbles of the gas through a very dilute solution of potassium bromide, pre- 
 pared by diluting a few drops of the laboratory solution with 5 cc. water. 
 Shake the solution with I cc. of carbon bisulfide. The brown color in the car- 
 bon bisulfide layer is characteristic of bromine. 
 
 12. Experiment. To 1 cc. potassium iodide solution add a drop of bromine 
 water. Shake with carbon bisulfide. The violet color is characteristic of iodine. 
 
 13. Questions. In each of the above experiments write the equation for the 
 reaction which has taken place. Point out the changes of valence and the inter- 
 change of electrical charges. Predict what reaction will occur when (a) chlorine 
 gas is passed through KI solution, (b) bromine is added to KC1 solution, and 
 (c) iodine is added to KBr solution. Experiment. Test the correctness 
 of your predictions. 
 
 14. The Conversion of Chemical into Electrical Energy. The reactions studied 
 in the present Assignment may be utilized to convert chemical into electrical 
 energy. In each of the reactions an interchange of electrical charges has taken 
 place; and the problem in the construction of an electrical cell is to arrange the 
 substances involved in the reaction in such a way that this interchange of 
 electricity can take place only through an external conductor such as a wire 
 joining the electrodes. Thus, in the reaction between zinc and copper ion two 
 changes of an electrical nature are involved: on the one hand zinc is changed 
 to zinc ion, and, on the other, copper ion is changed to copper. When the zinc 
 
 *The action of nitric acid on metals will be studied in Assignment 42. 
 
 [531 
 
is placed directly in a solution of a copper salt the electrical changes both occur 
 at the surface of the zinc and the transfer of electricity cannot be directly 
 detected; but when we arrange the materials so that the zinc electrode dips into 
 a zinc sulfate solution contained in a porous cup, and this is placed in a vessel 
 containing copper sulfate solution and the copper electrode, then no reaction 
 takes place. If now we join the two electrodes by a wire, electricity passes 
 through the wire and zinc will change to zinc ion at the surface of the zinc, 
 while copper ion changes to copper at the surface of the copper. The total 
 change in the cell is of course the same as when zinc acts directly on copper 
 ion. 
 
 15. Draw a diagram of the elecrical cell described above. Question. What 
 substances carry the electricity from one electrode to the other through the 
 solution ? 
 
 16. In a similar way the reaction between chlorine and bromide ion may 
 be used to generate electrical energy; but, to lead the current into and out of 
 the solution, we must use some substance, as graphite, which conducts elec- 
 tricity and is not attacked by chlorine or bromine. In one part of this elec- 
 trical cell there would thus be a graphite rod dipping into a solution which 
 contains bromide ion and bromine, and in the other part a graphite rod in a 
 solution of chloride ion and chlorine. The two solutions must be separated, 
 e. g., by a porous cup, in order to prevent the chlorine from reacting directly 
 with the bromide ion. When the graphite electrodes are joined by a wire a 
 current will flow, and at the one electrode bromide ion will change to bromine 
 while at the other electrode chlorine will change to chloride ion. 
 
 17. Problems. (/) What quantity of electricity (in coulombs) flows 
 through the circuit of the cell described in Paragraph 13 when one mol of zinc 
 disappears? What weight of copper will at the same time be deposited on the 
 copper electrode? 
 
 (2) Sketch the arrangement of a cell in which the reaction between 
 metallic copper and silver ion is used to produce an electric current. When a 
 current is taken from this cell what changes take place at each electrode? 
 
 (j) How many coulombs of electricity could be obtained by the transforma- 
 tion of 1 liter chlorine gas (measured at C and 760 mm.) to chloride ion? 
 
 ASSIGNMENT 42 
 
 OXIDATION OF METALS TO THEIR IONS. TABLE OF OXIDIZING 
 AND REDUCING AGENTS 
 
 Reference. Hildebrand, Chapter XV pages 226-245. 
 
 1. In Assignment 41 we found that certain metals can be oxidized by 
 hydrogen ion and that other metals cannot. We may say, therefore, that certain 
 metals are sufficiently strong reducing agents to reduce hydrogen ion to hydro- 
 gen and that others are not, and it follows that any reagent which will change 
 one of the latter metals to its ion is a more powerful oxidizing agent than is 
 hydrogen ion. In the present Assignment we shall study the action on zinc, 
 copper, and silver of two such oxidizing agents, nitric acid and bromine. 
 
 2. Arrange copper, silver, zinc and hydrogen, in a vertical column, in a 
 replacement series. Beside each symbol write the formula of the ion to which 
 the element can be oxidized. Your table now is a table of oxidizing 
 and reducing agents arranged according to their strengths, and this table 
 
 [541 
 
may be extended to include other substances than the metals and their 
 ions. Questions. Which is the more powerful reducing agent, Zn or Cu? Cu 
 or Ag? Which is the more powerful oxidizing agent, Ag+ or Cu ++ ? Ag+ or H + ? 
 (State the experimental facts upon which you based your answers to the pre- 
 ceding questions.) In describing the reaction between zinc and dilute hydro- 
 chloric or sulfuric acid solution is it correct to say that metallic zinc has com- 
 bined with chloride ion or with sulfate ion? How can the solid salts be pre- 
 pared from the solutions obtained in such experiments? 
 
 3. Action of nitric acid on zinc, copper, and silver. Experiment. Treat 
 about l / grams of zinc with 2 .V HNO 3 in a small beaker, covered with a 
 watch-glass. If the reaction is slow warm the mixture. The main reaction 
 under these experimental conditions is 
 
 3Zn (solid) + 2NO 3 - + 8H+ = 3Zn + * + 2NO (gas) + 4H 2 O. 
 
 4. Repeat the experiment with about 0.5 g. copper, and with a small 
 piece of silver. The reactions are similar to the one between zinc and 
 nitric acid, and the change from Zn to Zn + + corresponds exactly to the change 
 from Cu to Cu ++ and from 2 Ag to 2 Ag + . Write the equations for the reactions. 
 Question. How is the speed of the reaction between nitric acid and the metals 
 affected by a change in temperature and by a change in the concentration of the 
 acid? 
 
 5. Action of bromine on zinc, copper and silver. The halogens, chlorine, 
 bromine and iodine, all have a similar action on metals. To illustrate this 
 action we use bromine because its aqueous solution is a convenient reagent. 
 The solution of chlorine in water does not keep well, while iodine is a solid 
 which is only very slightly soluble in water. Experiment. Treat a small piece 
 of zinc with bromine water. After a few minutes pour the solution into a 
 porcelain dish, and heat it to boiling to expel any bromine that is present. To 
 a portion of the solution add a few drops of silver nitrate: AgBr, like AgCl 
 and Agl, is a difficulty soluble salt. Test another portion of the solution for 
 zinc ion. What evidence have you that a reaction takes place between zinc 
 and bromine? Repeat the experiment with copper and bromine water. Silver, 
 also, reacts with bromine. 
 
 6. The equation for the reaction between zinc and bromine solution is 
 
 Zn (solid) + Br 2 = Zn ++ +2Br 
 
 Since in this reaction zinc is oxidized to zinc ion we must conclude that bromine 
 is reduced when it changes to bromide ion. Write the equation for the action 
 of bromine on copper and on silver. 
 
 7. Review the replacement series for the halogens (Assignment 41) and 
 arrange chloride ion chlorine, bromide ion bromine, iodide ion iodine in a 
 table so as to give the relative strengths of the halogens as oxidizing agents and 
 of the halide ions as reducing agents. Add this table, in its proper place, to the 
 one previously prepared for the metals and their ions. Place Ca Ca + * and 
 Na Na + in this table. Show your completed table at once to the instructor. 
 
 8. We can determine the position of nitric acid in this table of oxidizing and 
 reducing agents by reference to the results of the experiment in Paragraph 6, 
 Assignment 34, when hot concentrated nitric acid was used to remove chloride ion 
 from a solution. Chlorine was formed and we can therefore decide that hot con- 
 centrated nitric acid is a stronger oxidizing agent than chlorine. Question. Is con- 
 centrated nitric acid a stronger oxidizing agent than bromine or iodine? 
 
 9. Study the method of balancing oxidizing and reducing reactions. Hilde- 
 brand, pages 226-233. 
 
 10. Problems, (i) Summarize the behavior of Na, K, Ca, Zn, Cu,. and other 
 
 [55] 
 
metals discussed in the lectures by dividing the metals into the three following 
 classes : 
 
 I. Those which react readily with water. Hydrogen is evolved. 
 II. Those which react readily with HC1 or H 2 SO 4 solution, but not 
 
 with water. Hydrogen is evolved. 
 
 III. Those which dissolve readily in HNO 3 or Br, solution, but not in 
 HC1, H 2 SO 4 or water. H 2 is not evolved. 
 
 Write the equation for the action of HC1 or H 2 SO 4 solution on metallic sodium. 
 
 (2) In the experiment in Paragraph 11, Assignment 41, manganese dioxide, 
 MnO 2 , reacted with chloride ion in acid solution to form chlorine. Manganous 
 ion, Mn ++ , was formed. Write the equation for the reaction. Predict what will 
 happen when copper is treated with MnO 2 in the presence of dilute sulfuric 
 acid. 
 
 (j) Sketch the arrangement of a cell in which the reaction between zinc 
 and bromine is used to produce an electric current. What changes take place 
 at the electrodes? State how you could bring about the reversal of all the reac- 
 tions in this cell. 
 
 ASSIGNMENT 43 
 FERROUS AND FERRIC IONS 
 
 Reference, in this and later Assignments : A standard text on Inorganic Chem- 
 istry. Note. Read again Asignment 31, Paragraph 2, and Assignment 38. 
 
 1. In Assignment 43 we shall study the chemistry of iron as an example 
 of a metal which forms two series of compounds in which the valence 
 of the metal is different. You will find on the laboratory shelves solutions and 
 solid salts of ferrous and of ferric iron. Give the valences of the following: 
 sulfate ion; sulfate in ferrous sulfate, FeSO 4 ; Fe (metal); ferrous ion, Fe ++ ; 
 ferric ion, Fe +++ ; iron in a ferrous salt, and in a ferric salt. Write the formulas 
 of the chlorides, nitrates, hydroxides, and oxides of ferrous and of ferric iron. 
 
 2. Many double salts are found among iron compounds. 
 FeSO 4 '(NH 4 ) 2 SO 4 ' 6H 2 O, ferrous ammonium sulfate, and 
 
 Fe 2 (SO 4 ) 3 '(NH 4 ) 2 SO 4 ' 24H 2 O, ferric ammonium sulfate (ammonium 
 iron alum) are examples of double salts whose solutions dissociate into the ions 
 of the simple salts. These two salts may be used in studying the properties of 
 Fe ++ and Fe +++ , respectively. Give the formulas of : alum, potassium iron 
 alum, sodium iron alum. 
 
 3. The ferrocyanides and ferricyanides are examples of double salts 
 whose solutions contain complex ions at high concentrations. The formula, 
 Fe(CN) 2 4NaCN, for sodium ferrocyanide shows its relation to the simple 
 salts. It can be prepared by dissolving the difficultly soluble ferrous cyanide, 
 Fe(CN) 2 , in excess of NaC'N, but in a solution of the pure salt the concentra- 
 tions of Fe ++ and CN~ are extremely small. Accordingly the formula is usually 
 written Na 4 Fe(CN) 6 to show that the ferrocyanide radical Fe(CN) is the 
 negative constituent of the salt. Similarly Na 3 Fe(CN) 6 is the formula 
 usually written for sodium ferricyanide. Questions. What is the valence 
 of the ferrocyanide radical? Of the ferricyanide radical? What is the 
 formula of potassium silver cyanide? How is it prepared? What ions are 
 present in its solution? Caution. Soluble cyanides and hydrocyanic acid are 
 deadly poisons : do not try unnecessary experiments with them. Ferro and 
 ferricyanides are not dangerous. 
 
 [561 
 
4. Color tests for Fe ++ and Fe + + + . Experiment. To a dilute solution of 
 FeCl 3 add a drop or two of K 4 Fe(CN) 6 solution. Repeat with K 3 Fe(CN) 8 . 
 Prepare a solution of a ferrous salt by dissolving a small amount of FeSO 4 
 or of FeSO 4 *(NH 4 ) 2 SO 4 in a little dilute sulfuric acid, and test portions of this 
 solution at once with K 4 Fe(CN) c and with K 3 Fe(CN) 6 . The character- 
 istic dark blue precipitates, or colloidal solutions, obtained in two of these 
 four experiments are probably identical. Equations may be omitted for these 
 reactions. 
 
 5. Oxidation of Fe ++ to Fe +++ . Experiment. Test for Fe +++ an acid solu- 
 tion of FeSO 4 which has been standing In the laboratory for several days. 
 The oxygen of the air slowly oxidizes Fe ++ to Fe +++ 
 
 4Fe ++ + O 2 + 4H + = 4Fe +++ + 2H 2 O. 
 
 Experiment. To a few drops of a ferrous salt solution in a porcelain dish add 
 a little HNO 3 , heat the mixture, dilute with water, and test whether the solu- 
 tion now contains Fe" 1 " 1 " or Fe +++ . Try an experiment with another oxidizing 
 agent and determine whether or not it oxidizes Fe ++ to Fe ++ +. 
 
 6. Note. Write equations for all reactions. 
 
 7. Reaction between Fe and Fe +++ . Experiment. Place some iron wire 
 or iron filings in a very dilute solution of FeCl 3 , stir the mixture, and test 
 small portions of the solution from time to time for Fe ++ and Fe +++ . 
 
 8. Reactions of iron ivith H + , and ivith Cu^. Experiment. Try the action 
 of dilute HC1 or H 2 SO 4 on a small piece of iron; heat the mixture, if neces- 
 sary. Predict whether ferrous or ferric ion is formed, and test the correctness 
 of your answer. Experiment. Cover a small piece of iron with a solution of 
 a copper salt. 
 
 9. From the results of the above experiments, or from any other source of 
 information, determine approximately the position of Fe Fe ++ and of Fe ++ Fe +++ 
 in your table of oxidizing and reducing agents, Assignment 42, Paragraphs 
 2 and 7. It is not advisable to attempt to assign a definite place in one large 
 table to each of the oxidizing and reducing agents which you will study. Thus 
 in the case of Fe ++ Fe +++ it is sufficient to note that Fe ++ Fe +++ , I~ 1 2 and 
 Ag Ag + are in the same part of the table: Fe ++ , I~ and metallic silver are all 
 readily oxidized by powerful oxidizing agents, while Fe + + + , I 2 and Ag + are 
 easily reduced by powerful reducing agents. Questions. Which is the stronger 
 reducing agent, Fe" 1 " 1 " or Fe? Which is the stronger oxidizing agent, Fe* ++ 
 or Fe ++ ? 
 
 10. Reduction of Fe +++ by H 2 S in acid solution. Experiment. To a few 
 drops of FeCL solution add 1 cc. 6 N HC1 and about 20 cc. water, and saturate 
 the solution with H 2 S gas. Boil the mixture in a porcelain dish to coagulate 
 the precipitated sulfur and expel the H 2 S, filter and at once test a portion of the 
 clear solution for Fe +++ . Treat the remainder of the solution again with H 2 S if 
 Fe +++ seems to be present. Name another reducing agent which might be 
 expected to reduce Fe +++ to Fe ++ and try the experiment. 
 
 11. Other reactions of Fe +++ . Experiments. Treat small quantities of dilute 
 solutions of FeCl 3 with the following reagents: 
 
 (a) NH 4 OH in small amount and in excess. The precipitate is Fe(OH) s . 
 
 (b) NaOH in small amount and in excess. 
 
 (c) Na 2 CO 3 or (NH 4 ),CO 3 . Filter; acidify the filtrate and test it for Fe +++ 
 with K 4 Fe(CN) 6 ; wash with water the precipitate (obtained by adding 
 carbonate), and test whether it is a carbonate or hydroxide. 
 
 (rf) NH 4 OH or NaOH and then a sulfide as H 2 S or (NH 4 ) 2 S. The black 
 precipitate is mainly Fe 2 S 3 , but may contain some FeS and S. 
 
 [57] 
 
12. Other reactions of Fe ++ . Repeat the experiments in (11) with small 
 quantities of dilute solutions of ferrous salt. In (a) and (b) note what 
 happens when moist ferrous hydroxide is exposed to the air; the oxygen 
 reacts much more rapidly with moist Fe(OH) 2 than with Fe ++ in acid solution. 
 The brown compound formed in the test for nitrate ion, Assignment 31, Par- 
 agraph 13, is an addition compound of NO and ferrous ion, probably Fe(NO) ++ . 
 This compound is formed when NO gas is passed into a cold ferrous salt solution, 
 and when nitric acid, nitrous acid, or NO 2 is reduced by excess Fe ++ . Questions. 
 What happens when a solution which contains this addition compound is heated? 
 Is the rate of reduction of nitrate ion in dilute acid solution by Fe ++ a rapid or a 
 slow reaction? Experiment. Determine whether the reduction of nitrous acid, 
 HNO 2 , is a rapid or a slow reaction by adding to 10 cc. water, 2 cc. FeSO 4 
 solution, a few drops oi 6 N H 2 SO 4 , and a drop of a nitrite solution. 
 
 13. Summarize in a table for future reference the reactions of Fe ++ and 
 Fe +++ with the following reagents: a strong oxidizing agent, as hot moderately 
 concentrated HNO 3 ; a strong reducing agent, as zinc in acid solution; H 2 S 
 in 0.3 N acid solution, and the reagents used in Paragraphs 11 and 12. Below the 
 table give methods of dissolving the difficultly soluble salts of iron which you 
 have studied. 
 
 14. Problems. Try experiments when necessary. 
 
 (/) How would you transform 
 
 (a) ferric chloride to ferric sulfate; 
 
 (b) ferric sulfate to ferric chloride; 
 
 (c) ferrous sulfate to ferric sulfate; 
 
 (d) ferric sulfate to ferrous sulfate? 
 
 (2) Fe(OH) 2 is moderately soluble in a solution of an ammonium salt, 
 as NH 4 C1, and Fe(OH) 3 is not. Outline an experiment to demonstrate this. 
 Which is the less soluble in water, Fe(OH) 2 or Fe(OH) 3 ? 
 
 (j) Suggest a method of preparing Fe(OH) 3 and ZnS from a single portion 
 of a solution which contains Fe ++ , Zn ++ , H + and SO 4 ~~. 
 
 ASSIGNMENT 44 
 MERCUROUS AND MERCURIC IONS 
 
 1. In the present Assignment we shall study the chemistry of another metal 
 which forms two series of compounds. You will find on the laboratory shelves 
 solutions of mercurous nitrate, HgNO 8 ; and of mercuric nitrate, Hg(NO 3 ) 2 . 
 What is the valence of mercury in a mercurous salt; in a mercuric salt? Write 
 the formulas of the chlorides, sulfates and oxides of mercurous and of mer- 
 curic mercury. 
 
 2. Mercurous chloride and mercuric chloride. Experiment. To 1 cc. 
 N HgNO 3 add 20 cc. water and 1 cc. 6 N HC1. Repeat the experiment with 
 N Hg(NO 3 ) 2 , and give a method of distinguishing between soluble mercurous 
 and mercuric salts. Note. Mercuric chloride is one of the few examples of a 
 weak salt. The concentrations of Hg ++ and of Cl~ in its solution are very small. 
 
 3. Reduction of mercuric ion in stages. Tin forms two series of compounds 
 stannous and stannic, in which the valence of tin is -\- 2 and + 4 respectively. 
 Stannous ion, Sn ++ , is a good reducing agent since it is easily oxidized to Sn +++ *. 
 The chlorides of tin are both soluble in water; and the common stannous salt 
 laboratory reagent is SnCl 2 solution to which HC1 has been added to prevent 
 
 [581 
 
the precipitation of a basic salt. Experiment. To 1 cc. N Hg(NO 3 ) 2 and 
 20 cc. water add one drop SnCL solution. Divide the mixture into two portions 
 and to one portion add SnCl 2 until it is present in excess. Observe the differ- 
 ence in appearance of finely divided mercury and a large drop of the metal. 
 Suggest a method of testing a solution for the presence of mercuric mercury. 
 Try your method with a dilute solution of HgCl 2 . 
 
 4. Reaction between mercuric ion and mercury. Experiment. To 1 cc. 
 N Hg(NO 3 ) 2 add a drop of mercury, shake the mixture gently, and after 
 several minutes add 20 cc. water. Pour off the solution and test it for Hg + 
 with HC1. Filter off the mercurous chloride and test the filtrate for mercuric 
 mercury. If there is more than a mere trace repeat the experiment and leave 
 the mercury in contact with the solution for a long time. The reaction be- 
 tween Hg and Hg ++ is similar to that between Fe and Fe +++ . In each case, in 
 the presence of excess of the metal, the reaction will proceed until at equi- 
 librium the concentration of the ion of higher valence, Hg ++ or Fe +++ , is ex- 
 tremely small. Questions. Which is the stronger oxidizing agent, Kg** or 
 Hg + ? Which is the stronger reducing agent, Hg + or Hg? 
 
 5. Reduction of Hg + and oxidation of Hg. By experiments with some of 
 the metals you have studied and small amounts of HgNO 3 solution determine 
 approximately the position of Hg Hg* in your table of oxidizing and reduc- 
 ing agents. In these experiments any metallic mercury formed can usually be 
 seen on the surface of the metal, but the solution may also be tested for the 
 dissolved metal after precipitating the mercurous mercury with HC1. From 
 your result predict whether metallic mercury will dissolve in dilute sulfuric 
 acid , and test your prediction. Experiment. Treat a very small drop of mer- 
 cury with about 30 cc. 2 N HNO 3 in a porcelain dish or beaker. If the reac- 
 tion is slow heat the mixture. While the reaction is in progress pour a small 
 amount of the solution into 10 cc. water and test for Hg + and Hg ++ . When 
 a moderate amount of the metal has dissolved, pour off the solution for use 
 in the next experiment. 
 
 6. Oxidation of Hg + and of HgCl. Experiment. Continue the preceding 
 experiment to determine how rapidly hot 2 N HNCX oxidizes Hg + . Repeat 
 the experiment with hot concentrated HNO 3 and a small quantity of HgNO 3 . 
 Experiment. Collect some HgCl on a filter. Treat a small portion in a porce- 
 lain dish with hot 2 N HNO 3 , and another portion with hot 2 N HNO 3 , to 
 which some HC1 has been added. Repeat the experiments, if necessary, with 
 more concentrated HNO 3 , and note whether the speed of oxidation is increased 
 by the presence of HC1. Experiment. Treat a small portion of HgCl with 
 bromine water, and test the solution as before for mercuric mercury. 
 
 7. Difficulty soluble salts of mercuric mercury. Experiment. Test the 
 action of each of the following reagents on a dilute solution of Hg(NO 3 ) 2 . Save 
 the precipitate for use in (8). 
 
 (a) H 2 SO 4 or a soluble sulfate. 
 
 (b) Na^COg. The precipitate is a basic carbonate. 
 
 (c) NaOH. The precipitate is HgO. 
 
 (d) NH 4 OH. When a chloride is present the precipitate is Hg(NH 2 )Cl, 
 ammonobasic mercuric chloride, which is closely related to Hg(OH)Cl, 
 basic mercuric chloride. Analogous but more complex substances are pre- 
 cipitated in the absence of chloride. What is the valence of the OH radical? 
 
 Of the NH 2 radical? 
 
 (e) H 2 S in- acid solution. 
 
 8. Predict the action of dilute HC1 or HNO 3 on each of the precipitates just 
 obtained (except HgS), and test your answers by experiments. 
 
 [59] 
 
9. Experiment. Collect some HgS on a filter and wash it with water. 
 Transfer part of the precipitate to a porcelain dish, add 2 TV HNO 3 boil the 
 mixture for about a minute, and filter. Test the filtrate for Hg*" 1 ". Repeat the 
 experiment with hot 2 N HNO 3 to which some HC1 has been added. From the 
 fact that HgQ 2 is a weak salt suggest an explanation for the result of this 
 experiment. Repeat the experiment with bromine solution instead of HNO 3 . 
 
 10. Other reactions of mercurous ion. Experiment. Treat a dilute solution 
 of HgNO 3 with each of the following reagents : 
 
 (a) H 2 SO 4 or a soluble sulfate. 
 
 (b) Na 2 C0 3 . 
 
 (c) NaOH. The precipitate is Hg 2 O, which, on heating, decomposes into Hg 
 and HgO. 
 
 (d) NH 4 OH. The precipitate is a mixture of metallic mercury and an am- 
 monobasic mercuric salt; see 7 (d). 
 
 (e) H 2 S in acid solution. The precipitate is a mixture of Hg and HgS. 
 
 11. From the results of your experiments in Paragraphs 6, 8, and 9, give 
 methods of dissolving each of the precipitates obtained in 10. Try experiments 
 whenever you are doubtful of your method. 
 
 12. Table of reactions for mercurous ion and mercuric ion. Summarize in 
 a table for future reference the reactions of Hg* and Hg ++ with the following 
 reagents: HC1; a strong oxidizing agent as Br 2 a strong reducing agent as 
 SnCl 2 (in small amount and in excess), and the reagents used in Paragraphs 7 
 and 10. Below this table classify the precipitates according to reagents used in 
 dissolving them and point out the considerations that led to the choice of each 
 reagent. Be sure that you can write the equation for each reaction. 
 
 13. Problems. Try experiments when necessary. 
 
 (/) Complete the following equations, and indicate what you would observe 
 in the corresponding experiments: 
 
 Sn+ + + HgCl 2 (excess) = 
 Sn + + (excess) + HgCl 2 = 
 
 (2) What reactions are possible between a sulfide and hydrochloric acid? 
 Between a sulfide and nitric acid? In connection with your answers discuss 
 the following experimental facts; CuS dissolves readily in hot 2 N HNO 2 but 
 not in hot 2 N HC1. 
 
 (j) Compare the action of hot 2 A r HNO 3 on CuS (Assignment 37 Paragraph 
 10) and HgS, and state what conclusion can be drawn with regard to the relative 
 solubilities of these two sulfides in water. Outline a method of preparing pure 
 HgCl 2 and pure CuSO 4 from a mixture of the freshly precipitated sulfides. 
 
 (4) Suggest methods of making the following preparations: 
 
 (a) HgO fromHg(N0 3 ) 2 ; 
 
 (b) Hg(NO 3 ) 2 from HgSO 4 ; 
 
 (c) Hg(N0 3 ) 2 fromHgN0 3 ; 
 
 (d) HgNO 3 from Hg(NO 3 ) 2 . 
 
 Note. It is recommended that Assignment 52 on Qualitative Analysis, follow 
 Assignment 44. 
 
 60 
 
ASSIGNMENT 45 
 LEAD ION, CHROMATE ION 
 
 1. In the present Assignment and the following ones we shall continue the 
 study of important common elements. Only very brief laboratory directions 
 will be given; each student is expected to plan the details of the experiments, 
 and is recommended to try additional experiments which are suggested by the 
 lectures or laboratory work. 
 
 2. Chromate ion, CrO 4 ~~, and bichromate ion, Cr,O 7 ~~. Note the colors 
 . of solutions of sodium (or potassium) chromate and bichromate. What is 
 
 the color of CrO 4 ~~? Of Cr 2 O r ~~? What is the valence of chromium in 
 these ions? Experiment. Add dilute HNO 3 or H 2 SO 4 to a small quantity 
 of a chromate solution. Add dilute NaOH or NH 4 OH to a bichromate so- 
 lution; acidify the solution obtained. Note. Many chromates are difficultly 
 soluble in water, but dissolve in HNO 3 . 
 
 3. Reactions of lead ion, Pb ++ . Experiment. Treat small portions of 
 0.1 A 7 Pb(NO 3 ) 2 with each of the reagents listed in Assignment 38 (Par- 
 agraph 5), with H 2 SO 4 , and with a soluble chromate. 
 
 4. Methods of. dissolving difficulty soluble lead salts. State which of the pre- 
 cipitates obtained in (3) you would expect to dissolve in 0.2 N HNO 3 . Experi- 
 ment. Prepare some PbS, collect it on a filter, and test portions of it with hot 
 2 N HC1, and with hot 2 N HNO 3 . Experiment. Collect some PbCl, on a 
 filter, pour cold water over it and test the filtrate for Pb* + , e. g. with H 2 S; 
 then pour hot water over PbCl 2 on a filter and test the filtrate for Pb ++ . 
 Experiment. Collect some PbCrO 4 on a filter, wash thoroughly with water 
 to remove any soluble chromate, and determine how readily it dissolves in hot 
 dilute HNO 3 . What ions are present in the solutions obtained? Experiment. 
 Collect some PbSO 4 on a filter, transfer some of it to a test tube, and pass in 
 H 2 S. Which is the less soluble, PbSO 4 or PbS? Experiment. Treat some 
 PbSO 4 with Na 2 CO 3 solution, as directed in Assignment 52, 11 (5). 
 
 5. Relative delicacy of tests for Pb + +. Experiment. Prepare 50 or 100 cc. 
 portions of 0.01 N, 0.001 N, 0.0001 N Pb(NO 3 ) 2 . Test 5 or 10 cc. of each 
 solution with the reagents used in (3). Specify the two most delicate tests 
 Name the two least soluble salts of lead. 
 
 6. Lead dioxide, PbO 2 . Experiment. Treat a small portion of PbO 2 with 
 hot 6 N HC1. What gas is evolved? Repeat the experiment with HNO 3 instead 
 of HC1. 
 
 7. Problems. (/) Outline methods for making the following transformations: 
 
 (a) PbSO 4 to PbS. (c) PbCO 3 to PbCrO 4 . 
 
 (b) PbS to PbSO 4 . (d) PbO 2 to Pb(NO 3 ) 2 . 
 
 (2) Name three amphoteric hydroxides, (cf. Assignment 35). What would 
 be observed on gradually acidifying a solution of sodium plumbite? 
 
 ASSIGNMENT 46 
 STANNOUS AND STANNIC IONS. AMPHOTERIC SULFIDES 
 
 Reference. Hildebrand, pages 197-199. 
 
 1. Ammonium polysulfide solution. Experiment. Prepare some colorless 
 ammonium sulfide solution by passing H 2 S into NH. t OH solution. Divide the 
 solution into three parts. To one portion add a very small amount of powdered 
 
 [61 1 
 
sulfur. Place another portion on a watch-glass to expose it to the action of the 
 oxygen of the air. Acidify the third portion with dilute HC1. Experiment. 
 Acidify small quantities of the laboratory solutions of ammonium polysulfide 
 (yellow ammonium sulfide), and of ammonium sulfide. Note. On account of 
 hydrolysis the (NH 4 ) 2 S solution is really NH 4 OH + NH 4 SH. 
 
 2. Reactions of Sn ++ and Sn ++++ . . Experiment. Prepare 0.1 N solutions 
 of stannous chloride and of stannic chloride, and determine the action of the 
 following reagents on each of these solutions : 
 
 (a) H 2 S in cold 0.3 N HC1 solution. These tests are delicate and characteristic. 
 
 (b) H 2 S in hot 2 N HC1. If there is a precipitate, filter, add water to the filtrate 
 and test it for tin. 
 
 (c) NH 4 OH in small amount and in excess. 
 
 (d) NaOH in small amount and in excess. If there is a precipitate after adding 
 NaOH in excess, filter, acidify the filtrate with HC1 or H 2 SO 4 and test it for 
 tin. 
 
 0) (NH 4 ) 2 S after NH 4 OH has been added. 
 
 (/) Ammonium polysulfide after NH 4 OH has been added. Pour through a filter 
 
 to remove any precipitate that may remain undissolved, and acidify with 
 
 dilute HC1. 
 
 (g) A soluble carbonate. Filter, wash the precipitate and test it for carbonate. 
 (h) A good oxidizing agent in dilute acid solution, such as bromine. 
 (i ) A good reducing agent in dilute acid solution, as zinc. 
 (/ ) HgCl 2 solution in small amount and in excess. 
 
 3. Action of concentrated HNO^ on Sn. Experiment. Treat a small piece 
 of tin with concentrated HNO 3 in a casserole. Evaporate the mixture just 
 to dryness: to avoid heating the residue strongly, evaporate the last portions 
 of the liquid by moving the casserole back and forth through a small flame. 
 Warm the residue with dilute HNO 3 , filter, and test the filtrate for tin. Note. 
 Stannic acid, or hydrated stannic oxide, when prepared in this way by the 
 action of concentrated HNO 3 on Sn or on a compound of tin, is difficultly 
 soluble in HNO 3 . Experiment. Suggest and try methods of dissolving this 
 substance. 
 
 4. Problems, (i) If a solution were known to be either SnCL or SnCl t how 
 would you identify it? 
 
 (2) How would you test for SnQ 2 in a solution of SnCl 4 ? 
 
 (j) What happens when HC1 is added gradually to a solution of (a) 
 sodium stannite, (b) sodium stannate, (c) sodium stannite after it has been 
 treated with a powerful oxidizing agent, as sodium peroxide, Na 2 O 2 , or hypo- 
 chlorite, (d) sodium sulfostannate? 
 
 (4) How would you prepare (a) ammonium polysulfide, (b) sodium sulfo- 
 stannate from SnS 2 , (c) sodium sulfostannate from SnS? 
 
 ASSIGNMENT 47 
 IONS OF ALUMINUM AND OF CHROMIUM. PEROXIDES 
 
 1. Write the formulas of aluminum chloride, nitrate and sulfate; chromium 
 chloride, nitrate and sulfate; alum and chrom alum. These salts are all readily 
 soluble in water and their solutions react acid to litmus. 
 
 2. Hydroxides. Experiment. To very dilute solutions of an aluminum 
 salt and of a chromium salt add NH 4 OH a few drops at a time until, after 
 shaking, the odor of ammonia is just perceptible. Warm the mixture. These 
 are delicate tests for Al +++ and Cr^*. 
 
 T621 
 
3. When a solution of an aluminum, or chromium, salt is treated with a 
 soluble carbonate or sulfide, the hydroxide of aluminum, or chromium, is pre- 
 cipitated. Explain. 
 
 4. Aluminate ion and chromite ion. A1O 2 ~ and CrCX" (or H.AICX') and 
 H,CrO 3 ~). Experiment. To dilute solutions of an aluminum salt and of a 
 chromium salt add a few drops of NaOH solution. To each mixture add 
 more NaOH and shake after each addition. Filter if a precipitate remains, 
 add some solid ammonium salt, and heat the mixture. 
 
 5. Ions of Chromium. State which ions of chromium can exist only in an 
 acid solution and which in an alkaline solution. What is the valence of 
 chromium in each of these ions? Suggest a method of changing bichromate 
 ion to chromic ion. 
 
 6. Reduction of Cr 2 O 7 ~~ and of CrO.,~~. Experiment. Test the action 
 of H 2 S on very dilute solutions of a bichromate in the presence of sulfuric 
 acid, heat the mixture to boiling, filter, and wash the precipitate. Make the 
 filtrate alkaline with NH 4 OH. Repeat the experiment with H 2 S and a very 
 dilute solution of a chromate in the presence of NaOH, heat, filter and wash 
 the precipitate. Acidify the filtrate. Experiment. Determine the action of hot 
 concentrated HC1 on Cr 2 O 7 ~~, and of Cl~ on CrO 4 ~~ in alkaline solution. 
 
 7. Oxidation of chromite ion. Experiment. Prepare a small quantity 
 of a chromite solution. To a portion of this add bromine solution; boil the 
 solution to expel any bromine that may remain. Question. How would you 
 prepare a hypochlorite solution? Write the equation for the oxidation of 
 CrO 2 ~ by C1O~. Experiment. Boil a small quantity of a chromic salt solution 
 with concentrated HNO 3 , or with aqua regia. 
 
 8. Peroxides. Sodium peroxide decomposes rapidly, even in a cold solution. 
 
 NiO HO = 2NaOH 
 
 Hydrogen peroxide in a cold, dilute acid solution decomposes only very 
 slowly. Na 2 O 2 is a very powerful and useful oxidizing agent. It reacts with 
 chromic hydroxide suspended in water, as follows : 
 
 2Cr(OH) 3 + 3Na 2 O 2 = 2CrO 4 " + 6Na + + 2OH~ +2H,O. 
 Experiment. Test the action of Na 2 O 2 on CrO 2 ~. The evolution of oxygen 
 observed in this experiment is due to the decomposition of some of the 
 peroxide. 
 
 9. In acid solution H 2 O 2 may be reduced by a powerful reducing agent, as 
 HoS; but it may also be oxidized by a powerful oxidizing agent, as Cr 2 O 7 ~~. 
 
 Cr 2 O 7 " + 3H 2 O 2 -f 8H + = 2 Cr +++ +3O 2 + 7H 2 O. 
 
 10. Problems. (/) Complete the following equations and state what would 
 be observed in the corresponding experiments: 
 
 Cr(OH) 3 + H+ (excess) = .................................................... 
 
 CrO 2 - + H + (excess) = .................................................... 
 
 CrO 4 - + H + (excess) = .................................................... 
 
 Cr 2 r ~ + H 2 S0 3 + H+ = Cr+" ............................................ 
 
 H 2 O 2 + MnO 2 + H+ = Mn" ............................................ 
 
 H 2 2 + I- + H+ = I, ................................................. 
 
 (2) State how you would prepare Cr 2 O r ~~ from Cr +++ , Cr(OH) 3 from PbCrO 4 . 
 
 (j) Concentrated NH 4 OH solution has a slight solvent action on Al(OH)g 
 and Cr(OH) 3 , but these hydroxides can be reprecipitated by adding an ammon- 
 ium salt and warming the mixture. Explain. 
 
 (4) List in a vertical column all the amphoteric hydroxides you have studied, 
 and write in two more columns the formulas of the corresponding ions which 
 
 [63] 
 
are formed when the hydroxide is dissolved in. excess of a solution of (a) a 
 strong acid, and (b) a strong base. 
 
 (5) Examine the position in the Periodic System (Hildebrand, page 257) of 
 each of the elements considered in the preceding problem, and write a brief 
 note on the gradations of properties in each family in which one of these 
 elements occurs. Name additional elements which might be expected to have 
 amphoteric hydroxides. 
 
 64 
 
SECTION V 
 QUALITATIVE ANALYSIS 
 
 ASSIGNMENT 51 
 
 THE DEVELOPMENT OF A SCHEME OF ANALYSIS FOR A LIMITED NUMBER OF 
 
 POSITIVE IONS 
 
 (To follow Assignment 38) 
 
 1. In Assignment 38 the chemistry of the positive ions already considered 
 was summarized, and the reactions of these ions with various reagents tabulated. 
 We shall now consider how this information may be used in analyzing an 
 unknown solution for these ions. 
 
 2. A possible method of analysis would be to devise a separate procedure 
 for each ion, then, if seven ions were to be considered, seven separate portions 
 of the original solution (or substance) would be taken, and each treated by an 
 experiment or series of experiments devised to show whether the particular 
 substance is present or absent in the unknown. Such tests were used in Assign- 
 ment 31 for Cl~, SO 4 ~~, and NO 3 ~; in these cases provision was made for 
 neutralizing OH~, because the presence of OH" would have interfered with or 
 spoiled the tests. It is to be noted also that the test used for Cl~ is not specific 
 for Cl~, since any other silver salt which is insoluble in dilute HNO 3 , as silver 
 iodide, would also be precipitated. It is evident that the application of this 
 method (of testing separate portions of the unknown for a particular substance) 
 becomes increasingly difficult as the number of substances which may be present 
 is increased. Also a lengthy procedure is often necessary, especially in testing 
 for certain positive ions when ions of several other metals are present. 
 
 3. Accordingly, in systematic qualitative analysis the method has been adopted 
 of using only one portion of the unknown solution in testing for metals, and 
 of preparing from it a pure compound of each metal. Such a process serves 
 to identify the positive constituents and also enables a comparison of the amounts 
 of the various metals to be easily made. The reagents ordinarily used in this 
 process are acids, NH 4 OH and NH 4 salts; and therefore separate portions of 
 the unknown are used in testing for H + , NH 4 + , and the various negative ions. 
 (The development of a systematic procedure for the detection of negative con- 
 stituents will not be considered in this course.) 
 
 4: As an illustration we shall discuss the preparation of pure salts of zinc and 
 calcium from an acid solution containing Zn ++ and Ca ++ . The Tables in Assign- 
 ment 38 show the following differences which may be utilized in devising methods 
 of separation : 
 
 (/) CaS is soluble in water and ZnS is difficultly soluble. The latter can 
 be precipitated, after the solution has been made alkaline with NH 4 OH, by add- 
 ing (NH 4 ) 2 S solution drop by drop, or better by passing in H 2 S gas. When 
 the ZnS has been completely precipitated, and removed by filtration, calcium 
 can be tested for in the filtrate and any desired salt of calcium can be pre- 
 pared. Experiment. Test this method by the Procedure given below, with an 
 acid solution which contains both Zn ++ and Ca ++ , prepared for example by adding 
 to 100 cc. water in a flask 2 or 3 drops of ZnSO 4 , CaCL and HC1 solutions. 
 Repeat, omitting the calcium salt, in order to determine if a satisfactory blank 
 is obtained in the calcium test. Procedure. (Be sure that you understand the 
 reason for each detail in the directions.) Add NH 4 OH, a few drops at a time, 
 until the mixture after shaking has a distinct odor of NH 3 , pass in H 2 S, shake 
 
 [651 
 
the mixture thoroughly, and cautiously note the odor to determine if H 2 S has 
 been added in excess; pass in more H 2 S if necessary. (If a large precipitate 
 is obtained, test with litmus and if the solution has become acid add more 
 NH 4 OH; again pass in H 2 S until it is present in excess.) Warm the mixture 
 to 50 or 60, and filter. Test the clear filtrate for calcium by adding (NH 4 ) 2 CO 3 
 solution, and heating the mixture almost to boiling. Note. If in the analysis 
 of an unknown solution no precipitate is observed in the test for Zn ++ or Ca ++ , 
 let the warm solution stand for several minutes before concluding that the ion 
 is absent. 
 
 (2) Zn ++ forms a complex ion with ammonia and Ca ++ does not. Both 
 carbonates are difficultly soluble in water, but ZnCO 3 , unlike ZnS, dissolves in 
 NH 4 OH. Accordingly CaCO 3 can be precipitated by treating the mixture with 
 NH 4 OH in excess and (NH 4 ) 2 CO 3 . Write out a Procedure (cf. preceding Par- 
 agraph) and note the reason for each of the experimental details. 
 
 (5) Zn ++ is converted into zincate ion by excess of a strong base. Pure 
 NaOH solution would precipitate some of the calcium as Ca(OH) 2 but 
 Na 2 CO 3 must also be added to make a satisfactory separation. The laboratory 
 NaOH solution contains some CO 3 ~~ but it is better to add Na 2 CO 3 . 
 
 5. Plan three methods of separating silver and calcium, i. e. of preparing 
 pure salts of silver and of calcium from a mixture of their nitrates. Write 
 out each method, stating why you think it will work; and show your note- 
 book to the instructor. 
 
 6. By combining your methods of separation, Paragraphs 4 and 5 devise 
 a "scheme of analysis" for these three positive constituents, silver, zinc and 
 calcium. Test your scheme by experiments with very small quantities of solu- 
 tions of (a) AgNO 3 alone, (b) ZnSO 4 alone, and (c) CaQ 2 alone. In each 
 experiment with a pure salt a satisfactory test should be obtained for one metal 
 and blanks in the tests for the other two. 
 
 7. Experiment. Predict what will happen when a small quantity of CuSO 4 
 solution is treated by your scheme of analysis, Paragraph 6, and test your answer 
 by an experiment. Devise a method of including copper in your scheme, and 
 test your method by experiments. 
 
 8. Assume that a sodium salt is present in solution with salts of silver, 
 copper, zinc and calcium, and consider how all the sodium originally present, 
 and no more, could be recovered as a pure salt. In other words, include sodium 
 in your scheme. Show your scheme of analysis for the five positive constit- 
 uents to the instructor. Question. With which of the five metals would 
 potassium have been found if it had been present in the original solution? - . 
 
 9. When an unknown solution has been analyzed in this way and a result 
 is inconclusive, owing for example to the small size of the precipitate, a charac- 
 teristic confirmatory test is made whenever possible, such as the dissolv- 
 ing of AgCl in NH 4 OH and the reprecipitation with excess HNO 3 , the flame 
 test for calcium, etc. 
 
 10. When the unknowns are solid salts an important part of the work is 
 the preparation of the solution for analysis. Preliminary experiments with small 
 quantities are performed to determine a method of dissolving the unknown, and 
 conclusions with regard to its nature can often be drawn from the results of 
 these experiments. The salts given out for analysis at this time will dissolve 
 either in water, in dilute acids or in concentrated acids. . A discussion of 
 methods of dissolving difficultly soluble salts will be given in Assignment 52. 
 
 11. Review your notes in the previous Assignments on the tests for 
 Na + , K + , NH 4 + , C1-, SO 4 ", CCX" and NO 3 -. Devise a method of testing for 
 
 [661 
 
S" based on the formation of H 2 S gas, and determine whether your test will 
 apply in the case of a very difficultly soluble sulfide, as CuS. 
 
 12. Analyses 1, 2 and 3. (Note. Do not begin these analyses until the 
 instructor has approved your scheme of analysis for the positive constituents, 
 Paragraphs 6 to 8. ) Test for the positive and negative constituents : 
 
 H, Ag, Cu, Zn, Ca, Na, K, NH 4 
 OH, Cl, N0 3 , S0 4 , C0 3 , S. 
 
 Record at once each experiment and observation and write down any con- 
 clusion that has bearing on the identification. Try to distinguish between large 
 amounts, small amounts, and traces of the different constituents. When you 
 report the result of an analysis to your laboratory instructor show him your 
 notebook record. 
 
 13. Problems. (/) What negative ions need not be tested for in solutions 
 which contain the. following positive ions: (a) H + . (b) Ca + *, (c) Zn ++ , (d) 
 Cu ++ , (e) Ag*? 
 
 (2) Point out the errors in the following reports: 
 
 (a) A solution known to contain Ag + was reported to contain SO 4 ~ ' 
 because BaCL gave a white precipitate which did not dissolve in 
 HN0 3 . 
 
 (b) A pure sodium salt which dissolved in water to give an alkaline 
 solution was reported to be a sulfate because a white precipitate 
 was obtained on the addition of Ba(NO 3 ) 2 - 
 
 (c) A pure salt which did not dissolve appreciably in water was reported 
 to be zinc nitrate. 
 
 ASSIGNMENT 52. 
 
 THE STANDARD SCHEME OF ANALYSIS. METHODS OF DISSOLVING DIFFICULTLY 
 
 SOLUBLE SUBSTANCES 
 
 (To follow Assignment 44) 
 
 1. In Assignment 51 a number of schemes were considered for the systematic 
 analysis of solutions containing Ag + , Cu" + , Zn ++ ,Ca ++ and Na + (and K + ). One of 
 the most satisfactory schemes, and the one used in practically all text books, 
 is the following: 
 
 To the solution add HC1 
 
 Precipitate 
 AgCl 
 
 Filtrate (concentration of H + , 0.3 A r ) : Saturate with H 2 S gas 
 
 Precipitate 
 CuS 
 
 Filtrate : Boil to expel H S, * add NH 4 OH and 
 (NH 4 ) 2 S. 
 
 Precipitate 
 ZnS 
 
 Filtrate: Add (NH 4 ),CO 3 and warm 
 
 Precipitate Filtrate : Evaporate to dry- 
 CaCO 2 ness, 
 
 [eat to expel NH 4 salts. 
 Residue: Na (and K) 
 salts. 
 
 * Note. It is really not necessary to expel the H 2 S ; but when* additional positive con- 
 stituents are included in the scheme, and several substances may be precipitated at this 
 point, the observation of the effect of adding NH 4 OH in the absence of sulfide often assists 
 in identifying the positive constituents. 
 
 [67] 
 
2. Iron and mercury in the scheme of analysis. From the results of your 
 experiments on the reactions of the ions of iron and mercury, Assignments 43 
 and 44, predict what would be observed if a solution containing each of the 
 following ions was treated by the scheme of analysis, Paragraph 1 : 
 Fe ++ , Fe +++ , Hg + , Hg ++ . Experiment. Verify your answers by experiments with 
 dilute solutions of each of these ions. Be sure that the blank tests for the other 
 positive constituents, e. g. Ca ++ , are negative. Question. What mistake has been 
 made if in the experiment with a mercuric salt a black precipitate is obtained 
 with sulfide in the NH 4 OH solution? 
 
 3. Separation of ZnS and FeS. Plan a method of dissolving a mixture of 
 ZnS and FeS and of obtaining pure Fe(OH) 3 and ZnS from the resulting solution 
 (which contains Zn ++ and Fe ++ ). Experiment. Test your method with small 
 amounts of (a) ZnS, (b) FeS, and (c) a mixture of ZnS and FeS. See Assign- 
 ment 51 Paragraph 4 (/) on the method of percipitating ZnS. 
 
 Separation of CuS and HgS. Experiment. Prepare a small amount, of (a) 
 CuS, and (b) a mixture of CuS and HgS. Treat each of these by the procedure 
 devised in Problem (j), Assignment 44. Questions. Why is the sulfur residue 
 black when CuS alone is used, (cf: Assignment 37, Paragraph 10)? Is a 
 confirmatory test for mercuric mercury necessary? For such a test see Assign- 
 ment 44, Paragraph 3. 
 
 Separation of AgCl and HgCl. Devise and confirm by experiment a method 
 of determining whether a chloride precipitate is either AgCl or HgCl. If a 
 large amount of HgCl is found and only a small amount of AgCl is present, 
 the method based on the use of NH 4 OH solution fails on account of the reduction 
 of Ag + by the metallic mercury. Devise a method of detecting silver in the 
 presence of a large amount of HgCl, based on the treatment of the mixture 
 obtained on the addition of NH 4 OH, with an oxidizing agent such as bromine 
 solution. Experiment. Try your method. 
 
 4. The other common positive elements are grouped as follows with the 
 elements already considered. 
 
 Division of Metals into Groups for Qualitative Analysis. 
 
 Characteristic Reagents for Precipita- Substances Precipitated. 
 
 Metal tion. 
 
 Name of Group. 
 
 I. Silver 
 
 Silver Group. 
 
 II. Copper 
 
 Copper Group. 
 
 [II. Zinc 
 
 Zinc Group. 
 
 IV. Calcium 
 
 Calcium Group. 
 
 HC1 or NH 4 C1 and an AgCl, HgCl, PbCl 2 par- 
 tially, BiOCl partially. 
 
 acid. 
 HS in 0.3 N acid. 
 
 NH 4 OH and (NH 4 ) 2 S 
 or NH 4 OH in excess 
 and H,S. 
 
 CuS, PbS, HgS, CdS, 
 Bi 2 S 3 , Sulfides of As, 
 Sb, Sn. 
 
 ZnS, A1(OH) 3 , Cr(OH) s 
 FeS, Fe 2 S 3 , MnS,NiS, 
 CoS. 
 
 (NH 4 ) 2 CO 3 andNH 4 OH. CaCO 3 , BaCO 3 , SrCO 3 , a 
 
 compound of rnag- 
 nesium when alcohol is 
 present, 
 none. 
 
 V. Sodium 
 
 Sodium Group. , 
 
 5. The further separation of the elements has already been illustrated in 
 Paragraphs 2 and 3, but only for two elements in each of three groups. The task 
 
 F681 
 
of developing a scheme of analysis becomes increasingly difficult as the number 
 of elements is increased, mainly on account of the large number of time-con- 
 suming experiments which must be made with known mixtures in order to 
 prove that the proposed methods are satisfactory. Therefore if you ever have 
 to make an accurate qualitative analysis of a complicated mixture it will be 
 necessary to consult a standard text-book, such as A. A. Noyes' Qualitative 
 Chemical Analysis, and to follow carefully the directions there given. Before 
 attempting to make an analysis with the aid of a book, work through the pro- 
 cedures with known solutions and study the chemistry involved. 
 
 6. The present course, however, is not primarily a course in Qualitative 
 Analysis, and a text-book will not be necessary. In the time which remains at 
 our disposal only a limited number of additional elements can be studied in the 
 laboratory. These will be investigated by the same methods as before, and 
 from the tables of reactions it will be possible to develop satisfactory methods 
 for the analysis of simple mixtures. 
 
 7 . Analyses Nos. 4, 5 and 6. In addition to the positive and negative con- 
 stituents considered in Assignment 51, test for iron, mercurous mercury and 
 mercuric mercury. Do not begin these analyses until the instructor has 
 approved your methods of analysis, Paragraphs 2 and 3. Also study carefully 
 the following paragraphs. 
 
 8. Analysis of an unknown solid. Record the appearance of the solid. Treat 
 a small portion with water and heat the mixture in a porcelain dish or beaker 
 if the solid does not dissolve readily. Always test the solution with litmus paper. 
 If the solution is neutral, determine if anything has dissolved, either by slowly 
 evaporating a few drops of the clear solution on a watch-glass, or by quickly 
 evaporating a larger quantity of the solution in a porcelain dish. 
 
 9. If the solid does not dissolve readily in water, in order to find a satis- 
 factory method of dissolving it, try preliminary experiments, such as are sug- 
 gested below, with small quantities of the powdered substance. The problem 
 is much simplified if the solid is known to be a single pure substance; for in 
 this case, when a reagent is found to react with some of the solid, the re- 
 mainder can be made to react in the same way by continued treatment with 
 this reagent. 
 
 10. When a method of dissolving the solid has been found, dissolve about 
 1 gram, and set aside about two-thirds of the solution in order that portions 
 of it may later be used in testing for NH 4 + and the negative constituents, and in 
 making confirmatory tests. Treat the remainder by the scheme of analysis for 
 the positive constituents ; use beakers or flasks in this work, and do not throw 
 away a portion of a solution whenever you have a large volume of it. 
 
 11. Methods of dissolving a difficultly soluble, pure substance. Observations 
 in these preliminary experiments with small quantities of the solid may be of 
 great assistance in identifying the substance. (In the following, some substances 
 are mentioned which will not be encountered in Analyses 4, 5 and 6). 
 
 (/) Treatment with dilute acid. Use N HNO 3 , boil if necessary. Note 
 the odor of any gas evolved. The use of HC1 is avoided because AgCl or 
 HgCl might be formed. Similarly BaSO 4 or PbSO 4 might form if H 2 SO 4 
 were used. 
 
 (2) Treatment with hot concentrated HC1. If chlorine is evolved a very 
 powerful oxidizing agent is present, such as MnO 2 or PbO 2 . Concentrated HC1 
 dissolves AgCl and HgCl readily, with formation of complex negative ions, but 
 the chlorides precipitate again when water is added to dilute the acid. 
 
 (3) Treatment with hot concentrated HNO 3 , a powerful oxidizing agent. 
 The rapid formation of brown gases indicates the presence of a reducing 
 
 [691 
 
agent. There is a residue of sulfur after the action of nitric acid on a sulfide, 
 and the solution may contain SO 4 ~~. 
 
 (4) Treatment with either (a) a mixture of concentrated HC1 and HNO 3 
 (aqua regia) or (b) bromine water. A few substances, as HgS, are dissolved 
 more readily by these reagents than by concentrated HNO 3 (which is really 
 a more powerful oxidizing agent) on account of the formation of weak electrolytes 
 or complex ions. 
 
 Note. Before using for analysis a solution containing concentrated acid 
 evaporate it almost to dryness to remove the excess of acid, and add water. 
 Question. What happens when H 2 S is passed into concentrated HNO 3 or 
 aqua regia? 
 
 (5) Treatment with Na 2 CO 3 solution to form a carbonate. Use a con- 
 centrated solution of Na 2 CO 3 in a porcelain dish, cover the dish with a watch- 
 glass and boil for several minutes. Filter and retain the nitrate to test for the 
 negative constituent. Treat the residue with dilute HNO 3 , filter if necessary, 
 and analyze the solution for the positive constituent. Review the action of 
 Na.,CO 3 solution on CaSO 4 , Assignment 32; and, experiment, try the above 
 procedure with BaSO 4 . 
 
 (6) Special tests. For example, AgCl dissolves in NH 4 OH solution, while 
 HgCl turns black; AgCl, when treated with water and H 2 S, is transformed into 
 the black Ag 2 S, which then may be collected on a filter and dissolved in con- 
 centrated HI\ f O 3 . 
 
 12. When the unknown is a mixture of solid salts, the same method of pro- 
 cedure may be followed in order to determine a method of dissolving the solid 
 completely. However, it is often convenient to analyze separately the portions 
 dissolved in the different operations. For example the mixture might consist of 
 one salt which dissolves readily in water or dilute HNO 3 , and a second salt 
 which does not. If this method is followed, care should be taken to dissolve the 
 first salt completely before continuing the experiments with the residue. Note. 
 Frequently when a mixture of two salts, as AgNO 3 and NaCl, is treated with 
 water a reaction takes place. 
 
 L3. Problems, (i) Briefly outline experiments by which you could identify 
 each of the following: AgCl, HgNO 3 , CuO, HgS, ZnCl 2 , Fe(OH) 3 , CaCO 3 . 
 
 (2 In order to be able to interpret quickly observations made in preliminary 
 experiments on difficultly soluble substances, Paragraph 11, prepare a table, 
 listing in a vertical row all the difficultly soluble compounds, studied thus far, and 
 in a horizontal row the reagents, (a) dilute HNO 3 , (b) hot concentrated HC1, 
 (c) hot concentrated HNO 8 , (d) bromine water (or aqua regia). (c) NH 4 OH 
 solution : If the compound is insoluble in a given reagent mark its position in 
 the table with a cross, if soluble write in the formulas of the substances formed 
 when it goes into solution. 
 
 ASSIGNMENT 53 
 LEAD, TIN, ALUMINUM, CHROMIUM, AND BARIUM IN THE SCHEME OF ANALYSIS 
 
 (To follow Assignment 47) 
 Reference. Assignment 52. 
 
 1. Lead in the scheme of analysis. State what would be observed if a solution 
 of pure Pb(NO 3 ) 2 were treated by the scheme of analysis which you used in 
 Assignment 52. Sugest a method of determining whether a white precipitate 
 obtained with HC1 contains any PbCL. What .would be the result of treating 
 
 [70] 
 
a mixture of HgS, CuS and PbS with hot 2 N HNO 3 (cf. Assignment 52, Par- 
 agraph 3) ? Give two reactions which are characteristic of Pb ++ but not of Cu ++ , 
 and one reaction which is characteristic of Cu ++ but not of Pb ++ . Suggest a 
 method of separating lead from copper and mercury, i. e. of including lead in the 
 scheme of analysis. Experiment. Test your method with known solutions. 
 
 2. Tin in the scheme of analysis. State what would be observed if solutions 
 of (a) stannous tin, (b) stannic tin were treated by the scheme of analysis. In 
 which analytical group would stannous and stannic tin be precipitated? Suggest 
 two methods of separating tin from copper, mercury and lead. Experiment. 
 Try one of your methods with known solutions. Repeat the experiment in 
 Assignment 46, Paragraph 2 (/), using small amounts of dilute solutions of 
 (a) Cu(NO 3 ) 2 , (b) SnCL, and (c) a mixture of Cu(N(X) 2 and SnCU. 
 
 3. Aluminum and Chromium in the scheme of analysis. State what would 
 happen if separate solutions of Al 2 (SO 4 ) a , Cr 2 (SO 4 ) 8 and K 2 Cr.,O 7 were 
 treated by the scheme of analysis. In which analytical group are aluminum and 
 chromium precipitated, and what are the formulas of the substances obtained? 
 How would you prepare from a mixture of ZnS, FeS, A1(OH) ; < and Cr(OH) 3 
 an alkaline solution containing HZnO.r, CrO 2 - and AlO.r? What would be 
 the effect of adding Na 2 O 2 to this alkaline solution? Outline a method of 
 showing the presence of aluminum, chromium and zinc in the resulting solution. 
 Note. The alkaline solution which contains CrO 4 ~~ must not be acidfied until 
 the peroxide present has been decomposed by boiling the solution. Experiment. 
 Test your method with known solutions containing aluminum and chromium, 
 and chromium alone. Experiment. Try the procedure with the reagents. NaOH 
 and Na 2 O 2 , and note whether a satisfactory blank test is obtained for aluminum 
 and for zinc. 
 
 4. Barium in the scheme of analysis. Explain why you might expect barium 
 to be found with calcium in the scheme of analysis. Look up in a table of 
 solubilities, such as is given on the front cover page of Smith's text-book, the 
 solubilities in mols per liter of the hydroxides, carbonates, sulfates and chromates 
 of barium and calcium. Arrange these salts in one column in the order of their 
 solubilities in water. Devise two methods for the separation of calcium and 
 barium. 
 
 5. Procedure for the separation of barium and calcium, after precipitation as 
 carbonates. Experiment. Test the following procedure with a freshly prepared 
 precipitate of (a) BaCCX, and (b) CaCO. v To dissolve the carbonate pre- 
 cipitate pour repeatedly through the filter a 10 cc. portion of hot 3 N acetic 
 acid, and use a little more acetic acid if the precipitate does not dissolve com- 
 pletely. Add 20 cc. N NH 4 Ac, and heat the solution to boiling in a flask. 
 Measure out 15 cc. A 7 " Na 2 CrO 4 and add it, a little at a time, heating and shak- 
 ing the mixture after each addition. Finally heat the mixture at 90 - 100 
 for one or two minutes, and shake it at the same time. Filter, even if you 
 cannot see a precipitate, remove the filtrate and wash the precipitate with cold 
 water. (Pale yellow precipitate, presence of barium.) To confirm the presence 
 of barium, dissolve the precipitate by pouring hot HC1 through the filter, evap- 
 orating to a small volume, and applying the flame test. To test for calcium in 
 the filtrate,- make the solution alkaline with NH 4 OH, add (NH 4 ) 2 CO 3 , or 
 ammonium oxalate, and heat the mixture just to boiling. Note. The separation 
 depends on the fact that BaCrO 4 is much less soluble in water than is CaCrO 4 . 
 If the BaCrO 4 is not precipitated slowly in a hot acid solution, as directed, the 
 precipitate will be so finely divided that it will run through the filter. 
 
 6. Analyses Nos. 7, 8, 9, etc. Test for all the positive and negative con- 
 stituents you have studied. Verify your results by experiments with known 
 substances. ! 
 
 [71 1 
 
7. Problems. (/) Complete the table of the behavior of difficultly soluble 
 substances towards various reagents, Assignment 52 Problem (<?), by the 
 addition of the -difficultly soluble substances studied in Assignments 45-47 and 53. 
 
 (2) Point out the errors in the following reports: 
 
 (a) Mercury is reported to be present because a black residue remained 
 when the sulfides of the copper group were boiled with 2 N HNO 3 . 
 
 (b) Lead is reported to have been found in the silver group but not in 
 the copper group. 
 
 (c) In the analysis referred to in (b) a black precipitate was observed 
 in testing for the zinc group, although no test for iron was otbained. 
 What errors in the method of analysis would produce this result? 
 
 (d) Aluminum is reported to be present in the zinc group, but the note- 
 book shows the following records : on precipitating the zinc group 
 a greenish precipitate was obtained on the addition of NH 4 OH, 
 later a yellow solution resulted when Na^CX was added ; when this 
 solution was acidified there was evolution of gas and the yellow 
 color disappeared; on the addition of NH t OH a precipitate was 
 formed. 
 
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